CHLORINE, BROMINE, IODINE AND ASTATINE

26. CHLORINE, BROMINE, IODINE AND ASTATINE A.

J. D O W N S and

C.

J.

ADAMS

University of Oxford

1. I N T R O D U C T I O N 1.1. GENERAL ATOMIC PROPERTIES^

The properties peculiar to non-metals can be attributed more or less directly to the relatively large effective nuclear charge experienced by the valence electrons of the atom, a characteristic reflected in the high ionization potential and electron affinity and the relatively small size of such an atom. Because electrons in the same shell shield one another relatively inefficiently from the nucleus, the incidence of non-metallic properties is a periodic function of atomic number, increase of which in a given period is accompanied by the transition from metallic to non-metallic behaviour. The noble gases, with their unique qualities of electron localization, are the culmination of this development. However, it is the preceding group of typical elements—the halogens—which provides, in physical and chemical terms, the best defined and most homogeneous family of non-metals, described elsewhere1 as "the most perfect series we have". The magnitudes of the effective nuclear charges associated with the halogen atoms are indicated by the ionization potentials, which are only 1-4 eV short of those of the corre­ sponding noble gas atoms, and by the relatively small sizes of the atoms. Further, the electron configuration of the halogen atoms in their ground states, ns2np5, just one electron short of the corresponding noble gas configuration, causes the atoms to be unusually powerful electron-acceptors, with electron affinities higher than those known for any other atomic species. Inasmuch as that rather elusive quantity electronegativity can be treated as an atomic property, it may be defined in numerous ways, but two useful criteria are (i) the mean of the ionization potential and the electron affinity (Mulliken's scale)4 and (ii) the effective nuclear charge experienced by an electron at a distance equal to the covalent radius from the nucleus (the Allred-Rochow formulation)5. In either case, the combination of the atomic properties outlined clearly implies exceptionally high electronegativity values for the halogen atoms. i N. V. Sidgwick, The Chemical Elements and their Compounds, Vol. II, p. 1097. Clarendon Press, Oxford (1950). 2 A. G. Sharpe, Halogen Chemistry (ed. V. Gutmann), Vol. 1, p. 1. Academic Press (1967); see also Codata Bulletin, International Council of Scientific Unions, Committee on Data for Science and Technology (1970). 3 W. Finkelnburg and F. Stern, Phys. Rev. 77 (1950) 303; R. W. Kiser, /. Chem. Phys. 33 (1960) 1265; M. F. C. Ladd and W. H. Lee, / . Inorg. Nuclear Chem. 20 (1961) 163; E. H. Appelman, /. Amer. Chem. Soc. 83 (1961) 805. 4 R. S. Mulliken, / . Chem. Phys. 2 (1934) 782. 5 A. L. Allred and E. G. Rochow, / . Inorg. Nuclear Chem. 5 (1958) 264. 1107

1108

CHLORINE, BROMINE, IODINE AND ASTATINE: A. J. DOWNS AND C. J. ADAMS

These and other numerical properties are summarized in Table 1, while Fig. 1 depicts variations of ionization potential, electron affinity and electronegativity for the halogens and neighbouring atoms in the Periodic Table.

FIG. 1. Ionization potentials, electron affinities and electronegativities of halogen and neigh­ bouring atoms in the Periodic Table.

Formation of Halides Perhaps the dominant feature of the chemistry of the halogens is the facility with which their atoms acquire an electron to form either the uninegative ion X - or a single covalent bond -X. As the oxidation potentials of Table 1 and other data suggest, however, this uninegative oxidation state becomes progressively less stable with respect to the free element with increase of atomic number. With the exception of helium, neon and argon, all the elements in the Periodic Table form halides which range in type from ionic aggregates at one extreme to simple molecules at the other. Halides generally are among the most important and common compounds, having played a central role in the historical develop­ ment of synthetic, structural and interpretative aspects of chemistry. In their capacity as donors, the halogen atoms can be treated formally, like hydrogen and alkyl groups, as one-electron ligands; in common with hydrogen and simple organic groups, e.g. -CH3, the halogens can also function as bridges between two other atoms, as in (SbF5)tt and Al2Br6. Some properties relevant to the formation of halides, including the covalent and ionic radii of the halogens, are shown in Table 1.

GENERAL ATOMIC PROPERTIES

1109

TABLE 1. SOME ATOMIC PROPERTIES OF THE HALOGENS2

Property Atomic number Electronic configuration First ionization potential (kcal) Electron affinity at 298°K (kcal) Electronegativity Dissociation energy of X2 molecule at 298°K, Z)(X2)(kcal) AÄ>°[X(g)] at 298°K (kcal) Single-bond covalent radius (A)* Ionic radius of X" ion (NaCl structure) (A) A/f/°[X-(g)] at 298°K (kcal) A#h°ydration[X-(g)]cat 298°K (kcal g ion"i) £ 0 (iX 2 /X"), aqueous solution (volts)

F

Cl

Br

17 9 35 [He]2s22p5 [Ne]3j23/>5 [Ar)3dms24p5 402

299

273

I

At a

53 85 [Kr]4di*5s25pS [Xe]4/H5
(-220)

810 40

84-8 30

790 2-8

72-1 2-5

(-71) (-2-4)

37-7 18-88

580 28-989

461 26-73

361 25-517

(-27-7) (-24)

0-71

0-99

114

1-33

(-1-4)

1-33

1-82

1-98

2-20

(-2-3)

-62-1

-55-7

-52-3

-46-6

(~ - 4 7 )

-121

-88

-80

-70

(~ - 6 6 )

+2-87

+ 1-356

+ 1065

+0-535

~+0-3

a Values for astatine given in parentheses are mostly estimated by extrapolation or related calculation (see, for example, ref. 3). b Given by one-half of the interatomic distance in the gaseous X2 molecules. c Absolute values relative to a hydration enthalpy for the proton of —260-7±2-5 kcal g ion - 1 .

Elementary Halogens The diatomic nature of the elementary halogens affords clear evidence of their almost unequivocally non-metallic behaviour: of the other elements, only hydrogen, nitrogen and oxygen share the property of being diatomic under normal conditions. Variations of such properties of the elementary halogens as their melting and boiling points and volatility are compatible with increases in magnitude of the van der Waals' forces with advancing size and polarizability of the atoms or molecules. Nevertheless, that intermolecular forces are significant in the condensed phases is shown (i) by the decrease in band gap for the solid elements with increasing atomic number (thus, with a band gap of 1 -3 eV6, iodine is com­ parable with various forms of phosphorus, arsenic and selenium), (ii) by the corresponding decrease in the ratio of the distance between next-nearest neighbours to that between nearest neighbours in the elementary solids (1 -68 for Cl 2 ,1 -46 for Br2 and 1 -29 for I 2 , see Table 11), (iii) by the degree of change of the vibrational spectrum of the diatomic molecule with the transition from the gas to the solid phase7, and (iv) by nuclear quadrupole resonance studies of the solid halogens8. The general trend of this evidence, together with the 6 N. N. Kuzin, A. A. Semerchan, L. F. Vereshchagin and L. N. Drozdova, Doklad. Akad. Nauk S.S.S.R. 147 (1962) 78; L. F. Vereshchagin and E. V. Zubova, Fiz. Tverd. Tela, 2 (1960) 2776; A. S. Balchan and H. G. Drickamer, / . Chem. Phys. 34 (1961) 1948. 7 M. Suzuki,T. Yokoyama and M. Ito,/. Chem. Phys. 50 (1969) 3392; 51 (1969) 1929; J. E. Cahill and G. E. Leroi, ibid., p. 4514. 8 C. H. Townes and B. P. Dailey, / . Chem. Phys. 20 (1952) 35; R. Bersohn, / . Chem. Phys. 36 (1962) 3445; N. Nakamura and H. Chihara, / . Phys. Soc. Japan, 22 (1967) 201; E. A. C. Lücken, Nuclear Quadrupole Coupling Constants, Academic Press, London (1969).

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CHLORINE, BROMINE, IODINE AND ASTATINE: A. J. DOWNS AND C. J. ADAMS

increasing stability of positive oxidation states in the series F, Cl, Br, I, makes it all the more intriguing to determine the nature of elementary astatine, about which very little is yet known. The intermolecular interactions also imply an ambivalence on the part of the halogen molecules which thus have some capacity to act either as donor or acceptor species. The acceptor nature of the heavier halogen molecules of the type XY (where X and Y are the same or different halogens) is amply supported by the very extensive class of complexes formed between XY and a wide range of electron-donors; such complexes range from polyhalide ions like I3 - to such labile charge-transfer systems as Br2,C6H6. Formally, the halogen molecule is normally regarded as the acceptor partner, but the self-association of these molecules, both in the solid state and in the discrete dimers Br4 and I4 recently identi­ fied9 in the gas phase and in solution, suggests that the donor function is not negligible. Intermolecular interactions in these systems extend from those weak enough to be classified as van der Waals' forces, e.g. Br4, to those recognisable as unambiguous covalent bonds, e.g. I3, a transition that is unusually illuminating in its relevance to some general problems of chemical bonding. Halogens in Positive Oxidation States The overall enthalpy changes (in kcal) accompanying the conversion iX2(standard state) -> X + (g)

are: F, +421; Cl, +328; Br, +300; I, +266-5; while for the process X2(aq)^X+(aq)+X-(aq) equilibrium constants of 10 ~ 40 ,10 ~30 and 10 _ 2 1 have been calculated when X = Cl, Br and I, respectively10. Together with other evidence, these data provide some guide to the stability of halogen cations X + . The removal of one electron from the valence shell of a halogen atom is unlikely to be accompanied by a substantial decrease in size, and it is this, with its consequent effects on lattice and solvation energies, rather than the magnitude of the first ionization potentials, that probably determines the infrequency with which free halogen cations occur under normal chemical conditions2. Nevertheless, cationic species such as H 2 OX + may well participate as intermediates in chemical reactions, whereas, for the heavier halogens, species containing coordinated X + units, e.g. BrF^, ICI2 and [I(NC 5 H 5 ) 2 ] + , have now been characterized (see Section 4). In non-aqueous media such as sulphuric or fluorosulphonic acids, bromine and iodine give coloured solutions whose contents probably include the cations X 2 and X3 ; recent studies of these solutions find, contrary to previous suggestions, no evidence of the species X + . In addition to the limited range of cations, there are known, for the heavier halogens but not for fluorine, many derivatives in which a formal positive oxidation state of + 1 , + 3 , + 5 or + 7 is associated with the halogen atom. Such systems invariably include ligands with electronegativities equal to or greater than that of the central halogen atom, the range thereby being restricted to oxygen compounds (e.g. CIO4), and interhalogen systems 9 L. J. Andrews and R. M. Keefer, Adv. Inorg. Chem. Radiochem. 3 (1961) 91; R. S. Mulliken and W. B. Person, Molecular Complexes, p. 141, Wiley, New York (1969); A. A. Passchier, J. D . Christian and N. W. Gregory, / . Phys. Chem. 71 (1967) 937; A. A. Passchier and N. W. Gregory, / . Phys. Chem. 72 (1968) 2697; M. Tamres, W. K. Duerksen and J. M. Goodenow, ibid., p. 966; D. D . Eley, F. L. Isack and C. H. Rochester, / . Chem. Soc. (A) (1968) 1651. 10 J. Arotsky and M. C. R. Symons, Quart. Rev. Chem. Soc. 16 (1962) 282.

GENERAL ATOMIC PROPERTIES

Uli

(e.g. BrF5 and ICI4). The oxidizing nature of the various oxyanions of chlorine, bromine and iodine is evident from the oxidation state diagrams of Fig. 2 which show that under acidic conditions most of the oxidation potentials are close to that of the oxygen couple 0 2 /H 2 0. Other noteworthy features of this figure are the instability of the elementary —Acid solution aH+—1 —Alkaline solution aOH_—l

FIG. 2. Oxidation state diagrams for the halogens in aqueous solution.

halogens with respect to disproportionation in alkaline solution, the comparative stability of the IOJ ion, the strongly oxidizing character of the recently discovered perbromate ion11, and the relative instability of the +3 oxidation state for all three halogens. Stereochemistry and Bonding The stereochemistries of some typical compounds of chlorine, bromine and iodine, summarized in Table 2, manifest a diversity of structural chemistry which is in striking contrast to that of fluorine compounds. The difference between fluorine and the heavier halogens in this respect arises from the practical inaccessibility to fluorine of oxidation states greater than 0 and of coordination numbers greater than 2 in covalently bonded systems. The formal expansion of the valence shells possible in chlorine, bromine and iodine is not easily rationalized because of the ephemeral nature of those atomic states which the chemist commonly refers to as "valence" states. Nevertheless, atomic properties such as (i) increased size and polarizability, (ii) decreased electronegativity, (iii) smaller energy separations between atomic ns, np, (« + l)s and (n+l)p states, and (iv) the availability of vacant nd orbitals are all likely to subscribe to the increased range of oxidation states and chemical environments open to the heavier halogens. 11 E. H. Appelman, / . Amer. Chem. Soc. 90 (1968) 1900; M. H. Studier, ibid., p. 1901; G. K. Johnson, P. N. Smith, E. H. Appelman and W. N. Hubbard, Inorg. Chem. 9 (1970) 119; J. R. Brand and S. A. Bunck, / . Amer. Chem. Soc. 91 (1969) 6500.

1112

CHLORINE, BROMINE, IODINE AND ASTATINE: A. J. DOWNS AND C. J. ADAMS TABLE 2. STEREOCHEMISTRIES OF COMPOUNDS OF CHLORINE, BROMINE AND IODINE

Coordination number

Examples

Geometry

1

Diatomic unit

Cl2, Br 2 ,1 2 , BrCl, HC1

2

Linear Angular

I^,CUCl-,BrIClC102, ClOJ, BrF^

3

Trigonal pyramid T-shaped unit

ClOr, BrOf, IOf C1F3, BrF3, RIC12 (R = organic group)

4

Tetrahedral unit Square planar unit Trigonal bipyramid with vacant equatorial site

CIOJ, BrO^, IOJ, C1207, FCIO3 IC17,12C16

5

Square pyramid

CIF5, BrF5, IF5

6

Octahedral unit Distorted octahedron

H 5 I0 6 , OIF5 IF6 (?), ΓΟβ units, e.g. in NH4IO3

7

Pentagonal bipyramid

IF7

I02F2", IF4+

Most of the stereochemistries represented in Table 2 are consistent with the formal disposition around the central halogen atom of 4,5 or 6 valence electron-pairs at the vertices, respectively, of a tetrahedron, trigonal bipyramid or octahedron. IF7 apart, clearcut structural information about systems with seven such pairs is generally lacking. Stereochemical details are then explicable on the basis of such factors as (i) the number of σ-bonding and lone-pair electrons on the central atom, (ii) the presence of more than one kind of ligand and hence more than one kind of σ-bond pair, and (iii) the minimization of the repulsive energy of interaction between the different electron pairs12. Hence, the structures of species like C1F3, IF4 and ICI4 can be rationalized. Although ττ-bonding is likely also to be important in derivatives like CIO4, there is no evidence to suggest that this has a major influence on the essential geometry of the unit. The conventional explanation of the expansion of the valence shells of chlorine, bromine and iodine depends on the nd orbitale, which are capable, in principle, of functioning as acceptors of σ- or π-electron density. Historically, this argument derives from the classical concept of the two-centre, two-electron bond, which leads to bonding descriptions of systems such as C1F3 or IF5 based 011 the use of dsp* or d2sp$ hybrid orbitale by the central halogen atom. Despite the lack of meaningful, quantitative information about the valence states of atoms, the one-electron orbital ionization energies of the free halogen atoms depicted in Fig. 3 furnish a useful comparison. According to such data, the energy separa­ tions between the valence np states and the excited states nd, (n + l)s and (n + \)p range from a maximum of ~370 kcal for fluorine to a minimum of ~ 170 kcal for iodine. Significantly, at the atomic level the (n+l)s orbitale are appreciably more stable than the nd orbitals. Since for effective participation in any bonding scheme it is a fundamental requirement of atomic orbitals that they should be compatible not only in symmetry but also in energy, these large energy separations must cast doubt on the validity of a concept involving the 12 R. J. Gillespie and R. S. Nyholm, Quart. Rev. Chem. Soc. 11 (1957) 339; R. J. Gillespie, Angew. Chem. Internat. Edn. 6 (1967) 819.

1113

GENERAL ATOMIC PROPERTIES

promotion of electrons to nd, (n+l)s or (n + \)p states and the subsequent formation of hybrid orbitals. It is generally acknowledged, however, that the radial functions and energies of nd electrons are unusually sensitive both to the formal charge on the atom and to the number of such electrons1*, the result being that the nd orbitals do participate in bonding to some, Br

ci --3A-

-P

_ 3d -\4p

_4d

-6p,5d -6s

-i\3s

-5p

100

"3p\

-4p\

-2p -5s 200

-3s

-4s

300

100 Kcal 10 eV -2s 400

xl0~ cm"1 FIG. 3. One-electron orbital ionization energies for the halogen atoms.

as yet undetermined, extent. Equally, it is recognized that the relative involvement of orbitals cannot be precisely prescribed as dsp* or d2sp*, for bonding may result solely from •s-orbital or ^-orbital or d-orbital interactions, or from any combination of these. In general terms, the halogens are electron-rich systems deficient in potentially bonding orbitals. In this respect, there is an obvious analogy between the heavier halogens and noble gases, as illustrated by the resemblance of iodine and xenon in the isoelectronic pairs XeFe, IF5; Xe0 3 , IOJ; XeOjj-, IO| _ . There is a less obvious affinity to the elements like gold or mercury with full or nearly full complements of d electrons: as with the halogens, the chemistry of these elements is strongly coloured by the dearth of bonding orbitals. The structural similarity of derivatives such as ICI2 and HgG^ or ICI4 and AUCI4 suggests an underlying conformity of bonding type. In these, as in noble gas compounds and in species 13 D . P. Craig and E. A. Magnusson,/. Chem. Soc. (1956) 4895; K. A. R. Mitchell, Chem. Rev. 69 (1969) 157; C. J. Adams, A. J. Downs and S. Cradock, Ann. Rep. Chem. Soc. 65A (1968) 216, and references cited therein.

1114

CHLORINE, BROMINE, IODINE AND ASTATINE: A. J. DOWNS AND C. J. ADAMS

such as HF2, the bonding can be described, to afirstapproximation, in terms of multi-centre molecular orbitals formed exclusively from those s- or^-orbitals of the central atom that are energetically eligible, the influence of rf-orbitals being neglected in the first place. For example, according to the simple LCAO-MO treatment developed by Rundle14 and others15, the molecular orbitals of the I3 ion are represented by a linear combination of the 5/?orbitals of the iodine atoms considered to overlap in the direction of the I—I—I axis. The formation of these three-centre molecular orbitals, with the associated energy-level diagram, is illustrated in Fig. 4, which provides a schematic description of so-called three-centre, four-electron bonding14. Although the contributions of s- and d-orbitals and of ττ-interactions should be incorporated in a more sophisticated account of the bonding, the simplified scheme of Fig. 4 appears more realistic than the "full hybridization" theory which gives

Y

x

Y

"G0G0G0 -ΘΘ

ΘΘ I

*. ΘΘ ΘΘ ΘΘ

-B

FIG. 4. Formation and energies of three-centre, four-electron molecular orbitals in the I3" ion and related systems.

excessive weight to energetically inferior d,sorp functions of the atom. By way of contrast, Fig. 5 illustrates in a purely general way the formation of a single σ-bond between a halogen atom and a second atom having just one available orbital; it is evident that the ray-orbital of the halogen makes a relatively greater contribution to the bonding when the second atom is the more electronegative partner. That the extent of ^-hybridization is usually small (< 10%) is suggested by various measurements, notably of the hyperfine coupling tensors for the radical anions XY~ (X, Y = the same or different halogen atoms)8. The problems 14 R. E. Rundle, Survey of Progress in Chemistry, 1 (1963) 81; / . Amer. Chem. Soc. 85 (1963) 112. is E. E. Havinga and E. H. Wiebenga, Rec. Trav. Chim. 78 (1959) 724; E. H. Wiebenga, E. E. Havinga and K. H. Boswijk, Ado. Inorg. Chem. Radiochem. 3 (1961) 133; E. H. Wiebenga and D . Kracht, Inorg. Chem. 8 (1969) 738.

1115

GENERAL ATOMIC PROPERTIES

posed by the bonding in I3 illustrate well the sort of impasse that has been reached at the theoretical level, arising partly from limitations of calculated wave-functions and partly from the difficulty of correlating molecular with atomic properties. The subtlety of problems of molecular structure and stereochemistry is further emphasized by a quite different approach which alludes to the properties not only of the ground but also of excited electronic states. The dependence of the energy of an aggregate upon its geometry

Electronegative atom

Halogen

Electropositive atom

FIG. 5. Correlation diagram for bonds formed by halogens with more electropositive or more electronegative elements.

may be expressed in terms of perturbation theory. If the Hamiltonian operator is ex­ panded16·17 as a Taylor series in 5«, a symmetry coordinate for molecular deformation: H=

B0+H;Si+iH?lS*+...

application of perturbation theory yields for the energy of the system, E = Ε°+<ψο | Hi I *β>Α+ \κφο | Hu' I φο>-Ύ<ψ°^*

L^m>2]&2+· · ·

where the subscripts o and m refer to the ground and to an excited electronic state, respec­ tively. The term in St, which describes Jahn-Teller distortions, vanishes for a non-degenerate ground state. However, for a degenerate ground state, such as that of BrF5 in the form of a trigonal bipyramid18, the term is finite: this implies instability with respect to a deformation which lifts the degeneracy of the system (achieved for BrF5, for example, by the transforma­ tion trigonal bipyramid -> square pyramid). The term in Sf describes the force constant appropriate to St; it is evident that a small value of (Em—E0), coupled with a non-vanishing matrix element <ψο\Η\'|0m>, can lead to a low or even negative value of the force constant. In the case of a totally symmetric ground state (which is the general rule), <^o\H'{\^y is non-zero only if H\ contains the same symmetry species as $m\ the characteristics of the force field are then uniquely determined by the lowest-lying excited states16. In general, if (Em—E0)^,4 eV and a deformation mode contains the same irreducible representation as the product ψοψη, the force constant appropriate to the appointed deformation is negative, 16 R. F. W. Bader, Mol. Phys. 3 (1960) 137. " D. H. W. DenBoer, P. C. DenBoer and H. C. Longuet-Higgins, Mol Phys. 5 (1962) 387; B. J. Nichol­ son and H. C. Longuet-Higgins, ibid. 9 (1965) 461. is R. S. Berry, M. Tamres, C. J. Ballhausen and H. Johansen, Acta Chem. Scand. 22 (1968) 231.

1116

CHLORINE, BROMINE, IODINE AND ASTATINE! A. J. DOWNS AND C. J. ADAMS

and distortion of the aggregate becomes energetically favourable. The system is then said to suffer a "pseudo-Jahn-Teller effect"19, for which general symmetry rules have been evolved to predict molecular structure. According to this treatment, molecular structure is not prescribed exclusively by the electronic properties of the individual atoms; rather does it depend also on the symmetry properties and relative energies of the ground and neighbouring electronic states of the aggregate and on the force field operative over the system. These features cannot justifiably be excluded from any meaningful interpretation of the structural chemistry of species like the polyhalide ions and charge-transfer complexes, for which neither the bonding energy nor (Em—E0) terms are large, and the frameworks of which are unusually pliable with respect to deformation. It is also a general precept that the exact structural characteristics of a polyatomic aggregate depend on its environment. That is to say that structure is a function also of the phase in which the unit occurs, or, in the event of solution, of the nature of the solvent, or, in crystalline solids, of features such as the packing of the structural units or the properties of counter-ions. Whenever there are relatively small energy differences between various structural alternatives open to a given system, such constraints, peculiar to the condensed phases, appear to be capable of impressing upon a polyatomic system a structure different from that of the isolated unit. The structure-determining influence of environment is exemplified by numerous structural features of the interhalogens and related species. 1.2. I R R E G U L A R I T I E S OF B E H A V I O U R OF T H E H A L O G E N S

Although the halogens form a remarkably homologous series, certain discontinuities are evident. Principal among these is the exceptional nature offluorine,as revealed by the small size of the atom in combination and of the F - ion, by its high first ionization potential (and electronegativity), and by the relatively low dissociation energy of the F 2 molecule. As a result, there is a much greater divergence of properties betweenfluorineand chlorine than between the other pairs of neighbouring elements. As in other series of elements, less striking but significant differences are found between the third and fourth members of the Group. These are reflected, for example, in the relatively irregular increases in size of the atoms and ions, such that chlorine and bromine are relatively similar to each other but differ from fluorine on the one side and iodine on the other. Variations in properties like electron affinity, polarizability and standard potential of the couple |X 2 /X~ (aq) display analogous irregularities. Estimates of the properties of astatine imply3 that it closely resembles iodine. This kind of variation in properties arises from the irregular increments of nuclear charge of 8,18, 18 and 32 in the series F, Cl, Br, I and At, attended by the subtle variations in the shielding from the nucleus of the valence p by the core electrons. At a chemical level we find that iodine differs from chlorine and bromine (i) in its significantly greater capacity to engage in chemical environments with relatively high coordination numbers (6 or even 7), (ii) in the susceptibility of iodides to oxidation, implied by the standard oxidation potentials of Table 1 and by the standard free energies of formation of pure iodides, e.g. HI, and (iii) in the relatively high stability of derivatives of the +5 oxidation state, e.g. I0 3 - and IF5, with respect both to oxidation and to reduction. The growing importance of the + 5 oxidation state with increase of atomic number of the halogens corresponds to a general feature of the B-group elements, that is, the emergence of the 19 L. S. Bartell, / . Chem, Educ. 45 (1968) 754; R. G. Pearson, / . Amer. Chem. Soc. 91 (1969) 4947.

THERMODYNAMIC ASPECTS OF THE CHEMISTRY OF THE HALOGENS

1117

oxidation state N- 2, where iVis the group valency. The phenomenon is commonly attributed to the so-called "inert-pair" eifect, a term which refers to the formal resistance of the ns2 electrons to ionization or to participation in covalent bond formation; it is most pronounced for the heavier elements of Groups II-V which are distinguished both by relatively large separations of the ns and np energy levels of the valence shell and by relatively low energies of chemical combination. For the halogens the "inert-pair" effect plays a less prominent role. This is evident from the fact that the ratio I8/Ip (where I8 and Ip are the ionization potentials for the removal, respectively, of an s and p electron) varies neither widely nor systematically; further, the average energies of bonds between a particular halogen and a more electronegative element (e.g. another halogen or oxygen) do not show the steady decrease with atomic number characteristic of the more metallic elements of Groups II-V. It has been suggested by more than one authority20 that with respect to combination with a more electronegative element there is an alternation of properties in the halogen series such that chlorine resembles iodine whereas bromine shares some of the anomalies exhibited by arsenic and selenium, its horizontal neighbours in the Periodic Table. Arsenic and selenium are noteworthy for the abnormal oxidizing properties associated with their group valence states of 5 and 6, respectively; features such as the failure to isolate arsenic pentachloride and the high oxidation potential of the couple Se04~(aq)/SeC>3~(aq) give notice of this property. In the past, abortive attempts to prepare either oxides of bromine or perbromates were widely taken to signify an inherent instability of these compounds. With the successful characterization first of bromine oxides21 and, in 1968, of perbromates11, reports of the non-existence of these species are seen to be an exaggeration. Much of the mystique of the "alternation" effect, at least in relation to the halogens, has thus evaporated. Accordingly, rationalization of the supposed non-existence of perbromates based, for example, on the spatial properties and energies of the 4d orbitals of bromine22, must now be seen for vain attempts to explain negative experimental evidence. Nothing so devalues the currency of a chemical theory as the discovery that the fact it purports to explain is itself incorrect, and the history of perbromates is a serious indictment of the chemical habit of theorizing about the "non-existence" of compounds involving significant degrees of electron-delocalization2^. This is not to gainsay the individual behaviour of bromine in many respects; thus, in practical terms, the oxides of bromine are of very low chemical stability21, while the thermodynamic barrier to the formation of perbromates11 demands unusually energetic methods of oxidation. However, the exact nature of much of this individuality and the extent to which thermodynamic or kinetic factors are implicated have yet to be properly defined. 1.3. THERMODYNAMIC ASPECTS OF THE CHEMISTRY OF THE HALOGENS

Despite the shortcomings of theoretical accounts, many aspects of halogen chemistry lend themselves unusually well to thermodynamic treatment both at an interpretative and at a predictive level. Two recent accounts2»24 illustrate clearly the way in which a chemical problem can be treated quantitatively by thermodynamic methods, there being ideally three 20 See for example C. S. G. Phillips and R. J. P. Williams, Inorganic Chemistry, Vol. 1, pp. 431, 663. Clarendon Press, Oxford (1965). 21 M. Schmeisser and K. Brändle, Adv. Inorg. Chem. Radiochem. 5 (1963) 41. 22 D . S. Urch, / . Inorg. Nuclear Chem. 25 (1963) 771. 23 W. E. Dasent, Non-existent Compounds, Edward Arnold and Marcel Dekker (1965). 24 D. A. Johnson, Some Thermodynamic Aspects of Inorganic Chemistry, Cambridge (1968).

1118

CHLORINE, BROMINE, IODINE AND ASTATINE: A. J. DOWNS AND C. J. ADAMS

stages in the procedure24: (i) the statement of the problem in thermodynamic terms; (ii) the restatement of the problem thermodynamically in the one of many possible ways that is the most susceptible to theoretical study; (iii) the theoretical interpretation of the restated problem. Thermodynamic methods cannot of themselves supply explanations for chemical phenomena; rather is their role to provide a precise and self-consistent rephrasing of these phenomena in terms of certain atomic or molecular energy terms. Since most reactions of the halogens and their inorganic compounds are fast (the hydrolysis of some non-metal halides and the reduction of perchlorate ion being among the few noteworthy exceptions), kinetic considerations are seldom of major importance2. Further, to the extent that the ionic model provides a useful theoretical basis for numerous, simple derivatives of the halogens, e.g. the crystalline alkali-metal halides and the solvated halide ions, the theoretical inter­ pretation of the thermodynamic data is direct, convincing and often at least semi-quantitative. It is for covalently bonded systems that the theoretical response is less satisfactory. Schematically the reduction of a free halogen, X2, in its standard state either (i) to an aquated halide ion X - , or (ii) to a crystalline ionic solid in conjunction with, say, some univalent cation M + is represented in Fig. 6. The free energy change attending the con,X-(aq) H lX2(standard) Atomisation

m

el

x(g)

,

^

2

0 ^

^(g)

M+X-(s) x* / 4, Λ J \ Atomisation -.,, N M v(standard) , · · .· ·· M+(g) V6/ 4- Ionisation

FIG. 6. Schematic representation of the formation (i) of an aquated halide ion and (ii) of a crystalline ionic halide M + X " .

version of one atom of the halogen in its standard state to the aquated halide ion under standard conditions of unit activity is seen to be equal to the balance of the free energies (a) of atomization, (b) of electron addition to the gaseous atom, and (c) of solvation of the gaseous anion. Since the entropy contributions are in all cases small compared with the enthalpy terms, and since moreover the entropy changes do not vary greatly from halogen to halogen in a comparative treatment of the problem, the free energy changes can be meaningfully approximated by the corresponding enthalpy terms. On this basis, for example, the standard redox potential E° for the couple £X2/X~(aq) is given approximately by E° = ~

F

[A#aX2(standard)-> X-(aq))-A#(H+(aq) -> *H2(standard))]

= ~i^H,°ms)]-EA~RT+AH^ r

+ 105]

(1)

L

where AHf°[X(g)] is the atomization energy of the halogen, EA the electron affinity, and A//aq° the enthalpy of hydration of the X - ion (all in kcal), Fis the Faraday and R the gas constant. Analogous arguments lead to the following expression for the standard enthalpy of formation of the ionic solid M+X - , Δ//}°, variations of which determine, to a good approxi­ mation, variations of AG/, that is, of the thermodynamic stability of the solid with respect

THERMODYNAMIC ASPECTS OF THE CHEMISTRY OF THE HALOGENS

1119

to its elements: AHf° = Atf/IXfe)] -EA + IM- Ü29S-2RT

(2) +

Here IM is thefirstionization potential of M and U29s the lattice energy of M X - at 298°K. Common to both processes is the conversion *X 2 (standard)->X-(g)

for which the standard enthalpies are given in Table 1; the total range of values in the series F, Cl, Br, I is less than 16 kcal, sufficiently narrow to suggest that variations in this quantity play a relatively minor part in differentiating the chemical properties of the halogens. The other two variables in equations (1) and (2) determined by the nature of the halogen are A//aq° and C/2985 both of which are inverse functions of the size of the ion. Thus, for a singly charged ion, AHaq0 is given approximately by the basically empirical relationship25 167

ΔΗΛ<10 =

faff

kcal g ion-i

(3)

wherein retu the "effective" radius of the ion, can be equated with the crystal radius of the halide ion. The lattice energy of the ionic crystal M +X - at 0° is related very nearly by the expression2 U0 =

NAeZ 7

r

p1 1 - -

(4)

to the electronic charge e, the Avogadro number N, the Madelung constant appropriate to the lattice-type A, the equilibrium interionic separation given, according to the simple model, by ΓΜ+ +Γχ- = r0, and the compressibility of the ions as determined by the (nearly constant) parameter />. It is the sensitivity of both A/faq0 and UQ (and the corresponding free energy terms) to variations in the size of the halide ion that accounts for the greatest part of the thermodynamic differences between the halogens with respect (i) to their oxidizing properties under aqueous conditions, and (ii) to the formation of crystalline, ionic halides. Thus, the data of Table 1 reveal variations in the values of AHm° of about 50 kcal g ion - 1 , while for the sodium halides the values of U0 (in kcal mol - 1 ) are:24 NaF, 216; NaCl, 185; NaBr, 176; Nal, 168. In the formation of a purely ionic halide of a metal in any given oxidation state, therefore, if we discount the small effect of possible change of structural type on lattice energy, the fluoride will always have the most negative ΔΗ/°. It can also be shown2»24 that for the reaction M» + X-n(s)+£X 2 (standard) -> M ( » + D + X - n + i ( s )

the enthalpy term, which is determined primarily by the difference in lattice energies, Δ{70> of the two ionic solids Mn+X-n and Μ ( Λ + 1 ) + Χ- Λ + 1 , becomes increasingly endothermic in the sequence X = F, Cl, Br, I. This is most clearly apparent on the basis of Kapustinskii's approximate expression26 for the lattice energy of an ionic crystal containing v ions per chemical formula with charges z+ and z_ and radii r+ and r_: ^ = ^^kcalmol-i

(5)

r++rn+

Hence, if the radii of the M

(w +1) +

and M

ions are taken to be approximately equal, being

25 Ref. 20, p. 163. 26 A. F. Kapustinskii, Quart. Rev. Chem. Soc. 10 (1956) 283.

1120

CHLORINE, BROMINE, IODINE AND ASTATINE: A. J. DOWNS AND C. J. ADAMS

denoted by r+, the difference in the lattice energies of MX w+ i and MXn is given roughly by Δ£/ο =

256(/i+2)(n+ l)-256(w+ \)n r++r x 512(n+l) kcalmol -1

It follows that the lattice energy of MXn+\ is invariably greater than that of MXW, the disparity decreasing as the size of the anion is augmented. Thus the variations in both — A///[X - (g)] and Δ£/0 place the capacities to stabilize the oxidation state n+1 in the order F > Cl > Br > I, manifest in the fact that many metals exhibit higher ionic oxidation states in fluorides than in other halides. Conversely, stabilization of the lower halide MX» with respect to oxidation to MX n + 1 is best achieved when X = I, consistent with the fact that low oxidation states in iodides are well-known features of the chemistry of many metals, e.g. iron(II), copper(I) and dipositive lanthanides. Arguments analogous to these serve also to account2»24 (i) for variations in the efficiency of metal halides as halogen-exchange reagents in reactions such as R-Cl+M + F-(s) -> R-F+M + Cl-(s)

and (ii) for the influence of cation size on the stability of crystalline derivatives of complex halide ions like HC12 ~ and IC12 ~, and for the mode of decomposition of such ions. The chemistry of molecular halogen compounds cannot be treated in this way because of the inaccessibility of quantitative information about changes such as (i) M"+(g) + nX-(g) or (ii) M (standard)

MXn(g) - M(g,valence state)

·- M(g, ground state)

• MXn(g) n/2 X2(standard) -

nX(g)

Nevertheless, certain helpful correlations can be made concerning the standard enthalpy of formation AH/0 of a gaseous molecular halide MXn in terms of the heats of atomization of M and X 2 and the average bond energy of MXW, 2?(M-X) (Fig. 7). A^°[MX n (g)] =

AHf0[M(g)]+nAHf°\X(g)]-nB(M-X)

(6)

If it is assumed that the enthalpy change associated with any condensation of gaseous MX» is small and that variations in TAS° terms from halogen to halogen are small compared with M (standard)

§X2 (standard)

AH?[M(g)] M(g)

ΔΗ?

-MX n (g) i

ηΔ·Η?[Χ(β)] nX(g)

-*B

FIG. 7. Thermodynamic cycle for the formation of a gaseous halide MX n from its elements.

SUBSEQUENT TREATMENT OF CHLORINE, BROMINE, IODINE AND ASTATINE

1121

those in ΔΗ0, variations in &Hf°[MXn(g)] provide a safe guide to those in AGf°[MXn]. Bond energy data (see Section 3) show that in all but a very few cases, the order of bond energies in any halide MXn is F > Cl > Br > I and that the variation is considerable. It follows from this and from the variations in A///°[X(g)] displayed in Table 1 that the stability of the halide MXn, with respect to its constituent elements, decreases in the order MFn > MC\n > MBrw > MIW. The pre-eminent thermodynamic stability of molecular fluorides is a function of the strong bonds that fluorine forms with other elements and the weak bond that it forms with itself; for the other halogens A/7/[X(g)] varies but little, and the thermodynamic stability of MXn is primarily a function of 2?(M-X). The relative stabilities of the halides NX3, OX2 and HX and the displacement of a heavier by a lighter halogen from a compound MXn illustrate this general thermodynamic condition. However, it is difficult to employ the same approach to interpret the relative capacities of the halogens to bring out high oxidation states because of the lack of definite information about the variation of the bond energy of MXn as n is varied. At best, the interpretative aspect of the bond energy treatment compares unfavourably with that involving the ionic model because bond energies are not theoretically comprehensible in the way that lattice energies are. But such theoretical shortcomings do not impair the value of the thermodynamic correlations, for the light that they may shed on the interdependence of molecular and atomic energy terms. 1.4. S U B S E Q U E N T T R E A T M E N T OF C H L O R I N E , B R O M I N E , IODINE AND ASTATINE

In Sections 2-4 chlorine, bromine and iodine are to be treated side by side in order to underline the homogeneity of much of their chemistry, and to delineate in more detail the resemblance and diversity of these three elements. The arrangement of the sections is illustrated schematically in Fig. 8. In the interests of uniformity within the framework of Comprehensive Inorganic Chemistry, the division of the sections is based on the formal oxidation state of the halogen, the elements (oxidation state 0) being treated first, followed Section 3

Section 2 Oxidation

Reduction Halide ions, Cl , Br , r and related species

+ e"

The elements Cl2, Br2, I 2

' ' + Donoΐ + Acceptor

Oxidation state -1

Section 4

Charge-transfer and related complexes

Oxidation state O

Derivatives of positive oxidation states: -ne~ A Cations B Oxygen compounds C Interhalogen compounds and polyhalide ions + Donor D Organic derivatives

Oxidation state > O

FIG. 8. Organization of the material in Sections 2-4.

1122

CHLORINE, BROMINE, IODINE AND ASTATINE: A. J. DOWNS AND C. J. ADAMS

initially by the halide ions and derivatives of the halogens with more electropositive elements (oxidation state — 1), and then by derivatives with more electronegative elements (oxidation state > 0). Regrettably, the scheme of things produces, of necessity, some artificial and arbitrary distinctions. For example, the classification according to oxidation state separates the diatomic halogens X 2 (Section 2) from the interhalogens such as C1F and IC1 (Section 4), whereas in reality the two classes are closely affiliated; the complication is particularly unfortunate for the apparent heterogeneity it brings to the treatment of donor-acceptor complexes of the two classes, whether these are molecular like Br2,QH6 or Me3N,ICl, described mainly in Section 2, or ionic as in the polyhalide species I3 - and IBr2 ~, presented with the interhalogens in Section 4. On the other hand, it is inappropriate to consider the heaviest known halogen, astatine, within the context of the scheme of Sections 2-4. Because of the short half-lives of even the most stable isotopes, macroscopic quantities of astatine cannot be accumulated. Accord­ ingly, our knowledge of its chemistry depends· entirely on tracer studies. The very limited state of this knowledge, combined with the practical aspects peculiar to its chemistry, requires that a separate section be devoted to astatine. 1.5. PSEUDOHALOGENS 27 · 28

Reference must also be made here to the class of pseudohalogens, first defined by Birckenbach and Kellermann27, comprising such molecules as cyanogen, (CN)2,oxocyanogen, (OCN)2, thiocyanogen, (SCN)2, selenocyanogen, (SeCN)2, and azidocarbondisulphide, (SCSN3)2. Like the diatomic halogen molecules, these contain two relatively electro­ negative units (X) directly linked, as in NC-CN; in common with the halogens, they are reduced characteristically to anions X~ or to molecular pseudohalides, e.g. CH3-X. Examples of pseudohalide anions are CN~, OCN~, SCN~, SeCN - and SCSN3-, as well as TeCN~ 29 and N 3 ~, for which the pseudohalogen parents are not known. The physical and chemical properties diagnostic of pseudohalogens or pseudohalide ions are, in outline: 1. The parent pseudohalogen is a molecular compound involving a symmetrical com­ bination of two radicals, X-X. With alkalis the reaction is often analogous to that of the halogens, e.g. (CN) 2 +20H- ^ C N - + O C N " + H 2 0

The pseudohalogens also undergo some of the addition reactions typical of halogens, as in CH2=CH2+(SCN)2 -> NCSC 2 H 4 SCN

2. The pseudohalogens react with various metals to give salts containing X - anions; the salts of silver, mercury(I) and lead(II) are typically sparingly soluble in water. 3. The hydrides HX are acids which are, however, very weak compared with the halogen acids, as illustrated by the following pKa values: HCl, - 7 - 4 ; HN 3 , 4 4 ; HCN, 8-9. The difference in acidities is presumably a function primarily of the differences (i) in H-X bond energy and (ii) in hydration energy of the X~ ion. 2 ? L. 28

Birckenbach and K. Kellermann, Chem. Ber. 58 (1925) 786, 2377. T. Moeller, Inorganic Chemistry, p. 463. Wiley (1952); R. C. Brasted, Comprehensive Inorganic Chemis· try (ed. M. C. Sneed, J. L. Maynard and R. C. Brasted), Vol. 3, p. 223. Van Nostrand (1954); ref. 20, p. 467; F. A. Cotton and G. Wilkinson, Advanced Inorganic Chemistry, 2nd edn., p. 560. Interscience (196Q R. B. Heslop and P. L. Robinson, Inorganic Chemistry, 3rd edn., p. 551. Elsevier (1967). 29 A. W. Downs, Chem. Comm. (1968) 1290.

PSEUDOHALOGENS

1123

4. The pseudohalogen radicals form compounds among themselves, e.g. NC-SCN and with the halogens, e.g. C1N3; species analogous to the polyhalide ions are also known, e.g. (SeCN) 3 - and [I(SCN)2] - 30. 5. The pseudohalide ions form complexes with metals as do the halide ions, e.g. [Co(NCS)4]2 -, though the stabilities of analogous halide and pseudohalide complexes differ widely in many cases. 6. The pseudohalogen forms molecular compounds analogous to molecular halides, e.g. CH3NCO and Si(NCO)4, though, in contrast with the halogens, the asymmetry of the pseudohalogen radical may lead to isomeric forms depending on the mode of coordination of the radical, as in CH3-SCN and CH3-NCS. 7. In common with halide ions, a pseudohalide ion is oxidized to the parent pseudo­ halogen by suitable oxidizing agents. A typical reaction is 2Fe3 + +2SCN" -> 2Fe2 + + (SCN) 2

More detailed aspects of pseudohalogen chemistry are referred to in the context of carbon, nitrogen and appropriate Group VI elements and of transition metal complexes. Except in relation to derivatives of the halogens in positive oxidation states (e.g. interhalogens and polyhalide ions), pseudohalogens are not otherwise treated in this section. Although the classification of materials as pseudohalogens has some practical value, albeit of a limited and largely empirical nature, the affinity of such materials to the halogens themselves must not be exaggerated; even the cursory outline given here suggests a number of significantly divergent properties. Further, it should be appreciated (i) that few of the classical pseudohalogens measure up completely to the idealized behaviour which has been described, and (ii) that, if such deviations are tolerable, then the class is capable of very considerable enlargement^1; for, in terms of at least some of their properties, species such as (N0 2 ) 2 , (RS)2 (where R is an organic group), (OH)2, [OS(F)02]2 and (C103)2 deserve also to be treated as pseudohalogens. GENERAL REFERENCES BRASTED, R. C , Comprehensive Inorganic Chemistry (ed. M. C. Sneed, J. L. Maynard and R. C. Brasted), Vol. 3, van Nostrand (1954). COTTON, F. A. and WILKINSON, G., Advanced Inorganic Chemistry, 2nd edn., Interscience (1966). Gmelins Handbuch der Anorganischen Chemie, 8 Auflage: "Chlor", System-nummer 6, Teil A and B, Verlag Chemie, Weinheim/Bergstr. (1968-9); "Brom", System-nummer 7, Verlag Chemie, Berlin (1931); "Iod", System-nummer 8, Verlag Chemie, Berlin (1933). GUTMANN, V. (ed.), Halogen Chemistry, Vols. 1-3, Academic Press (1967). HESLOP, R. B. and ROBINSON, P. L., Inorganic Chemistry, 3rd edn., Elsevier (1967). JOLLES, Z. E. (ed.), Bromine and Its Compounds, Benn, London (1966). MELLOR, J. W., A Comprehensive Treatise on Inorganic and Theoretical Chemistry, Vol. II, Longmans, Green and Co., London (1922); Supplement to Mellor's Comprehensive Treatise on Inorganic and Theoretical Chemistry, Supplement II, Part I, Longmans, Green and Co., London (1956). PASCAL, P., Nouveau Traiti de Chimie Minerale, Vol. XVI, Masson et Cie, Paris (1960). PHILLIPS, C. S. G. and WILLIAMS, R. J. P., Inorganic Chemistry, Vol. 1, Clarendon Press, Oxford (1965). SIDGWICK, N. V., The Chemical Elements and Their Compounds, Vol. II, Clarendon Press, Oxford (1950). WELLS, A. F., Structural Inorganic Chemistry, 3rd edn., Clarendon Press, Oxford (1962).

30 C. Long and D. A. Skoog, Inorg. Chem. 5 (1966) 206. 3i See for example L. Birckenbach and K. Huttner, Chem. Ber. 62 (1929) 153; Z. anorg. Chem. 190 (1930)1.

1124

CHLORINE, BROMINE, IODINE AND ASTATINE: A. J. DOWNS AND C. J. ADAMS

2. THE ELEMENTS CHLORINE, BROMINE AND IODINE 2.1. DISCOVERY AND HISTORY3233

Of the halogens, chlorine was the first to be discovered, being prepared by Scheele in 1774 by heating hydrochloric (muriatic) acid with manganese dioxide34, though the fumes of the gas must have been known from the time of the thirteenth century by all those who made and used aqua regia. The preparation of hydrochloric acid itself (then called spiritus salis) was first reported in a fifteenth century Italian manuscript35, while sodium chloride, known as "salt" from the earliest times, is referred to by Pliny in his Naturalis Historiae^. In accordance with the views of Lavoisier that all acids contained oxygen, chlorine was first named oxymuriatic acid, a view apparently supported by observations of Berthollet (i) that, if the manganese dioxide is first deprived of some of its oxygen by calcination, it furnishes a smaller quantity of Scheele's gas, and (ii) that oxygen is evolved when the gas reacts with water. After attempts by Gay-Lussac, Thenard and Davy to decompose the so-called oxymuriatic acid had failed, it was Davy who in 1810 first proposed the elementary nature, the name chlorine (from the Greek χλωρός, green) and the symbol Cl for the gas37. At about the same time, the group name halogen (from the Greek α λς, sea-salt, and the root yev -, produce) was first coined to describe the salt-forming tendencies of the individual mem­ bers of the group. Next in order of discovery was iodine, first prepared in 1812 by the French chemist Courtois. When an aqueous extract of the calcined ashes of seaweed, that is, kelp, was treated with sulphuric acid, it was noted by Courtois that the black precipitate first produced was converted on heating to liberate the new element in the form of its violet vapour38. The name iode, the French equivalent of its present name (Greek ίώδης, violet), was designated in 1813 by Gay-Lussac, who also demonstrated some striking analogies between the new substance and chlorine, and in a famous communication "Memoire sur Fiode" (1814)39 described large tracts of iodine's chemical behaviour. Of the three halogens, bromine had the least eventful history, its elemental nature and its relation to chlorine and iodine being recognized from the very first. The discovery of the element is credited to Balard in 1826 in the course of studying the mother liquor remaining after the crystallization of salt from the water of the Montpellier salt-marshes, which is rich in magnesium bromide40. Balard was attracted by the intense yellow coloration which developed when chlorine water was added to the liquor; ether-extraction followed by treatment with potassium hydroxide destroyed the colour, while the residue was shown, 32 J. W. Mellor, A Comprehensive Treatise on Inorganic and Theoretical Chemistry, Vol. II, Longmans, Green and Co., London (1922); Supplement to Mellor's Comprehensive Treatise on Inorganic and Theoretical Chemistry, Supplement II, Part I, Longmans, Green and Co., London (1956). 33 Gmelins Handbuch der Anorganischen Chemie, 8 Auflage: "Chlor", System-nummer 6, Teil A, Verlag Chemie, Weinheim/Bergstr. (1968); "Brom", System-nummer 7, Verlag Chemie, Berlin (1931); "Iod", System-nummer 8, Verlag Chemie, Berlin (1933). R. C. Brasted, Comprehensive Inorganic Chemistry (ed. M. C. Sneed, J. L. Maynard and R. C. Brasted), Vol. 3, van Nostrand (1954). 34 C. W. Scheele, König. Vetens. Akad. Stockholm, 25(1774)89; Opuscula chimica et physica, Leipzig, 1 (1788) 232; The Chemical Essays ofC. W. Scheele, p. 52. London (1901); Alembic Club Reprints (1897) 13. 35 L. Reti, Chymia, 10 (1965) 11. 36 Pliny, Naturalis Historiae, Book 33, chapter 25 (first century A.D.). 37 H. Davy, Phil. Trans. 100 (1810) 231; Alembic Club Reprints (1894) 9. 38 B. Courtois, Ann. Chim. 88 (1) (1813) 304. 39 J. L. Gay Lussac, Ann. Chim. 91 (1) (1814) 5. 40 A. J. Balard, Ann. Chim. Phys. 32 (2) (1826) 337.

NATURAL OCCURRENCE

1125

when heated with manganese dioxide and sulphuric acid, to produce red fumes which condensed to a dark brown liquid with an unpleasant smell. There is no question but that the element had been isolated by Joss and by Liebig41 prior to Balard's discovery; however, neither of these investigators recognized the true nature of their product, Joss mistaking it for selenium and Liebig for iodine chloride. On the other hand, Balard was unquestionably the first to appreciate the elemental nature of the material and its relation to chlorine and iodine. The substance was first called muride, but the name bromine—from the Greek βρωμος, a stench—was later preferred. 2.2. N A T U R A L

OCCURRENCE323342-48

The halogens are found in nature, not in their highly reactive elemental states, but most commonly in the form of the corresponding halide anions; in the rather exceptional case of iodine, iodate deposits are an important natural source of the element. It has been estimated that igneous rocks, which constitute about 95% of the earth's crust, contain on average the following approximate proportions of the halogens in the combined state: Cl, 0-031%; Br, 1-6x10 _4 % ; I, 3 x 10 ~5%. In accordance with the similarity of radius of the chloride and bromide ions, the mineral chemistry of the two elements is quite closely related, and bromine is known to replace chlorine in numerous minerals. The bulk of the chlorine and bromine present in sedimentary and volcanic rocks takes the place of OH groups in hydroxide-bearing minerals such as hornblendes, micas, clay materials and aluminium hydroxide. There are otherwise known relatively few chlorides or bromides which may be considered as minerals in the strict meaning of the word; the relatively insoluble ores of certain heavy metals (e.g. AgCl, horn silver; AgBr, bromargyrite; Ag(Br,Cl), embolite; Hg2Cl2, calomel; CuCl, nantokite) are of negligible commercial importance. However, the largest natural source of chlorine and bromine is the sea; out of a total average salinity of about 3-4%, sea water contains approximately 1*9% chlorine, mainly as sodium chloride though with smaller amounts of other chlorides, 0.0065% bromine (representing a chlorine :bromine mass ratio of nearly 300:1) and 5 x l 0 - 8 % iodine. Isolated bodies of water in arid regions are frequently found to have a high chloride content; the Great Salt Lake of Utah, for example, contains no less than 23% sodium chloride, while the Dead Sea, with a total salinity of more than 30%, contains near the surface about 3-5% calcium chloride, 8-0% sodium chloride and 13-0% magnesium chloride. Whereas the chloride ion makes up 90% of the total anion content of sea water, the main anion of river water (average salinity 0-01-0-02%) is the bicarbonate ion, the chloride ion amounting to only 2-5%. Both the chloride and bromide ions from weathered rocks are thus dissolved 4i J. R. Joss, / . prakt. Chem. 1 (1834) 129; J. v o n Liebig, Liebig's Ann. 25 (1838) 29. 42 F . C. Phillips, Quart. Rev. Chem. Soc. 1 (1947) 9 1 ; D . T. Gibson, Quart. Rev. Chem. Soc. 3 (1949) 2 6 3 ; V. M . Goldschmidt, Geochemistry, Oxford (1954); V. V. Cherdyntsev, Abundance of Chemical Elements, Chicago University Press (1961); G. G a m o w , Biography of the Earth, Macmillan, London (1962); L. H . Ahrens, Distribution of the Elements in our Planet, McGraw-Hill, N e w York (1965). 43 C. S. G. Phillips and R. J. P. Williams, Inorganic Chemistry, Clarendon Press, Oxford (1965-6). 44 N . V. Sidgwick, The Chemical Elements and their Compounds, Vol. II, p. 1139. Clarendon Press, Oxford (1950). 45 Z. E. Jolles (ed.), Bromine and its Compounds, Benn, L o n d o n (1966). 46 Kirk-Othmer Encyclopedia of Chemical Technology, 2nd edn., Vols. 1, 3 and 11. Interscience (1963-6). 47 C. A . Hampel (ed.), The Encyclopedia of the Chemical Elements, Reinhold, N e w York (1968). 48 J. S. Sconce (ed.), Chlorine: its Manufacture, Properties and Uses (A. C . S. Monograph N o . 154), Reinhold, N e w York (1962).

1126

CHLORINE, BROMINE, IODINE AND ASTATINE: A. J. DOWNS AND C. J. ADAMS

and transported into the hydrosphere. They concentrate in the seas whence they may return to the lithosphere through deposition from parts of the sea separated from the main body of the ocean. The question of whether the primordial ocean contained the same salts in proportions that are the same as, or similar to, those of today's oceans is still a moot point45, though the discrepancy between the sodium xhlorine ratio in the ocean and that in rocks makes it difficult to subscribe to the view that the composition of the ocean has remained constant. To account, at least in part, for the balance in nature of chlorides and bromides, it has been suggested that carbonates have reacted, and may still be reacting, with hydrogen chloride and bromide present in volcanic gases. The principal sources of chlorine and bromine of commercial significance are as follows: Ocean-derived Deposits of Such Minerals as Rock Salt, NaCl; Sylvine, KCl; Carnallite, MgCl2, KCl, 6H20; Kainite, MgS04, KCl, 3H20 These deposits are widely distributed in the world, many of them being of immense propor­ tions; countries favoured by such natural resources include, for example, the United States, Russia, Germany, Austria, Italy, the United Kingdom and Brazil. For the production of chlorine, sodium chloride is the principal raw material, and where this is obtained from natural rock salt deposits, it may either be mined or pumped to the surface as a saturated aqueous solution. The bromine content of carnallite deposits is no more than 0-1-0-35%, while rock salt contains typically 0-005-0-04% bromine. Minerals richer in bromine, such as bischoffite and tackhydrite (MgCl2 ,6H 2 0 and CaCl2,2MgCl2,12H20, respectively) are not abundant and are less important as sources of bromine than is carnallite. Bromine is not extracted directly from carnallite but from the mother liquor remaining after the extraction of potassium chloride from the mineral. Natural Brines Derived from Seas and Lakes These are now a significant source of chlorine and the major source of bromine. Solar evaporation is commonly used to concentrate the brines and to isolate the principal dissolved ingredients. In certain areas, special topographical and climatical conditions greatly facilitate this process; such is the case in the Rann of Cutch in India, in the saline lagoons of the Black Sea and Caspian Sea, in the Sebkha-el-Melah in Tunisia, and on the shores of inland seas and lakes such as the Dead Sea and Great Salt Lake. Although the average bromine content of sea water is only 0-0065%, some isolated bodies of water are much richer in the element: e.g. the Dead Sea, 0-4-0-6%; Sebkha-el-Melah, ~ 0-25%; Sakskoe Ozero (Crimea), 0-28%; Searles Lake (California), 0-085%. At the present time, bromine is extracted commercially from normal ocean water containing no more than 0-0065% bromide ion. Wells and Springs Apart from the brine produced artificially by pumping water into underground salt deposits, natural brines from wells and springs are also very widespread; thus, natural waters associated with oil-fields are often relatively rich in halide ions. Although these are most commonly the result of leaching of salt layers, it is possible that some originate from underground pockets left by ocean concentrates. Many such wells and springs are found, for example, in the United States, mainly in Michigan, Ohio and West Virginia, in parts of Russia, Israel and Italy, and many provide important natural sources of common salt,

FORMATION OF THE ELEMENTARY HALOGENS

1127

calcium chloride, magnesium chloride and bromine. Whereas chloride is the principal anion of most of the waters, the bromide content varies somewhat unpredictably from zero to proportions (typically 0-1-0-4%) compatible with the commercial production of the element. Iodine differs in a number of respects from chlorine and bromine. In the first place, it is, by a considerable margin, the rarest of the three halogens. Like chlorine and bromine, it is widely distributed in nature in rocks, soils, underground brines and sea water, but always in low concentration (see Table 36). Unlike chlorine and bromine, iodine occurs naturally not only as iodide but also as iodate, and it is in this form (in the minerals lautarite, Ca(I0 3 ) 2 , and dietzeite, 7Ca(I03)2,8CaCr04) that iodine is found in the Caliche beds in Chile, which, with an iodine content of 0-02-1%, remain the major commercial source of iodine. The concentration of iodine in other secondary deposits is insignificant, while the iodide content of sea water is so low as to preclude serious commercial exploitation for the production of iodine. Apart from the Chilean nitrate deposits, some natural brines, commonly derived from oil-well drillings, also contain commercially useful concentrations of iodine (typically 30-40 ppm); such are the natural brines of Michigan and, now on a diminished scale, the oil-well brines of California. Despite the low concentration of iodine in sea water, some seaweeds, notably the deepsea varieties, can extract and accumulate the element. Those of the Laminaria family are the richest, containing up to 0-45% on a dry basis. The ashes of seaweed provided the first commercial source of iodine. Although largely superseded as the chief raw material for the production of iodine by the Chilean deposits and by natural brines, seaweed continues to be used locally to produce iodine, principally in Japan. A number of other types of marine life, such as oysters, sponges and certain fishes, also concentrate iodine in their systems, and iodides and iodates in sea water enter into the metabolic cycle of most marine flora and fauna. In the human body, the greatest concentration of iodine is found in the thyroid gland. Iodine appears to be a trace element essential to animal and vegetable life. Bromide ions are found in living organisms, but their biological role is unknown and no precise physio­ logical significance has so far been established. In plants and animals chlorine exists primarily as the chloride ion dissolved in cell liquids, but organo-chlorine compounds, mostly of fungal origin, are also found in nature (see pp. 1336-40). 2.3. FORMATION OF THE ELEMENTARY HALOGENS32»33·45"48

General Methods of Production The formation of free chlorine, bromine or iodine may be effected by one of two chemical methods: (i) oxidation of halide derivatives, and (ii) reduction of compounds of the halogens in positive oxidation states, e.g. ClO _ or IO3 ~. Of these two general methods, the first is very much more important than the second for such practical reasons as the ready availability and cheapness of the starting materials, the comparative facility of selective oxidation rather than selective reduction, and the kinetic inertness exhibited by some of the commoner "positive" halogen derivatives, e.g. CIO3- and C10 4 ~. The thermodynamic balance between the two processes is well illustrated by reference to the oxidation state diagrams of Fig. 2 and to the differences between the standard redox potentials of the couples

1128

CHLORINE, BROMINE, IODINE AND ASTATINE: A. J. DOWNS AND C. J. ADAMS

XO3-/4X2 and JX2/X~ on the one hand and of the couples Χ0 3 -/£Χ 2 and XO3-/Xon the other. It is thus evident that the majority of reagents capable of reducing the X0 3 ~ ion to the free halogen in acid solution are also capable of reducing it to the X~ ion. Oxidation of Halides Since the free energy of formation of halides, whether pure or in solution, decreases in the series chloride > bromide > iodide, the thermodynamic barrier to the production of free halogen from the corresponding halide decreases in the same order. For chlorine, therefore, the range of chemical agents that will accomplish this oxidation is limited to the more energetic oxidizing agents, whereas for iodine oxidation is easily accomplished and strongly oxidizing conditions are neither necessary nor yet desirable in that they may lead to "positive" iodine compounds. An example of chemical oxidation to produce chlorine is provided by the reaction of hydrogen chloride with manganese dioxide Mn02+4HC1 -> MnCl 2 +Cl 2 +2H 2 0

Apart from its historical significance34, this has for long been a method of producing chlorine on a small scale in the laboratory, and was also the basis of the Weldon process formerly used for the manufacture of chlorine. Other effective oxidizing agents of hydrogen chloride or metal chlorides include dichromates, permanganates, lead dioxide, sodium bismuthate, peroxydisulphates and chlorates. Air oxidizes hydrogen chloride at elevated temperatures and in the presence of certain catalysts, e.g. CuCl2, as in the Deacon Process (q.v.). Processes depending on the oxidation of chlorides either by nitric acid or by sulphur trioxide have also been developed on, or up to, commercial proportions. However, electrolytic oxidation of chloride solutions accounts for more than 99% of the chlorine of commerce32»33»46 _48 . In an analogous manner, bromine is liberated from hydrogen bromide, metal bromides or solutions of these by oxidation with reagents such as manganese dioxide, nitric acid or bromates. The Deacon Process of air-oxidation is applicable to the conversion of hydrogen bromide to bromine, as is electrochemical oxidation of bromide ions. But the only methods of importance for the manufacture of bromine are based on the oxidation of bromidecontaining solutions by chlorine32»33*45 ~47. Formation of iodine from iodides is accomplished by agents such as chlorine, bromine, manganese dioxide, concentrated sulphuric acid, nitric acid and iodates; even such mild oxidizing agents as Fe 3+ , [Fe(CN)6]3~, Cu 2+ and antimonate(V), under appropriate conditions of pH in aqueous solution, oxidize iodide ions quantitatively to free iodine. Because of the ease with which iodine can be estimated volumetrically, a number of these reactions, e.g. that with iodate and copper(II), afford noteworthy methods of quantitative analysis. For the production of iodine on the large scale, oxidation of iodides by manganese dioxide, chlorine, nitrite or dichromate has been exploited. Electrolytic oxidation of iodide solutions provides another potential route to iodine, though it is commercially unrealistic. Conversion of iodides to iodine in commercial quantities is now mostly accomplished by chlorine oxidation in processes that are similar in principle to those used for the extraction of bromine32»33»46»47. A modification of detail, though not of principle, is provided by the thermal decom­ position of certain metal halides, whereby the elementary halogen is released, though the method is of little practical consequence as a preparative procedure. In thermal

FORMATION OF THE ELEMENTARY HALOGENS

1129

decompositions such as PtCl4 ->PtCl2+Cl2 PtCl2 ->Pt+Cl2 PbCl4 ->PbCl2+Cl2 2AuBr3 ->2Au+3Br2 U0 2 Br 2 -*U0 2 +Br 2 2EuI3 -^2EuI 2+I2

the halides range from molecular to predominantly ionic species; the molecular compounds are generally characterized by relatively weak metal-halogen bonds, whereas for the ionic systems decomposition is favoured (i) by the presence of small, highly charged cations, and (ii) by unusually large ionization potentials governing the conversion of the less to the more highly charged form of the metal. Reduction of Derivatives of the Halogens in Positive Oxidation States Perhaps the most important reaction of this class is that between halate and the corre­ sponding halide ions in acid solution: 5 X + X 0 3 +6H + ^3X 2 +3H 2 0

In all cases the balance of this reaction is highly sensitive to pH, and under alkaline conditions the halogens are themselves unstable with respect to disproportionation into XO3" and X ~ ions. Thus, in one process for the manufacture of bromine from sea water, bromine vapour is absorbed in sodium carbonate solution when it forms bromide and bromate45'46; subsequent acidification regenerates free bromine in more concentrated form. Similarly, the naturally occurring iodates found in the Chilean nitrate deposits are converted to iodine by the following processes46: I0 3 " + 3HS03" -> I" + 3HS04" 5 I + I 0 3 + 6 H + ->3I 2 +3H 2 0

The reaction between chlorate and hydrochloric acid has also been used for the laboratory preparation of chlorine32»49. The oxidation state diagrams of Fig. 2 make it clear that, despite the ready dispropor­ tionation of the free halogens under alkaline conditions, the interaction not only of a halate but also of a hypohalite, a chlorite or perhalate with the corresponding halide under acidic conditions is thermodynamically capable of producing the free halogen. The reactions and

HOC1+C1- + H + -► C1 2 +H 2 0 IO4- +71- +8H + -> 4I 2 +4H 2 0

afford examples of the practical working of this principle, though the thermodynamic potential of some systems is not realized because of kinetic factors (associated, for example, with the perchlorate ion), or because of side-reactions (e.g. disproportionation). Reaction of both oxyhalogen and interhalogen derivatives with other relatively mild reducing agents 49 E. Dufilho, Bull. Soc. Pharm. Bordeaux, 63 (1925) 41.

1130

CHLORINE, BROMINE, IODINE AND ASTATINE: A. J. DOWNS AND C. J. ADAMS

may also lead to the elementary halogen, as in the following examples: I2O5 + 5CO 2C103- +I 2 5(COOH)2+2I03- +2H + 6H5I06+15H2S 2NH3+2C1F3 6Nb + 10BrF3 6CIO2+2BrF3

aqsoln aqsoln aqsola

> l2+5C02 > 2IO3- +C12 > 10CO2 +1 2 + 6Η2Θ > 3I 2 +2H 2 S0 4 +13S+28H 2 0 ^N 2 +6HF+C1 2 > 6NbF5 + 5Br2 > 6CIO2F+Br2

Again, thermal decomposition of many oxyhalogen and polyhalide derivatives produces the free halogen. Thus, one mode of decomposition of aqueous hypobromous acid proceeds slowly at room temperature according to the equation: 4HOBr

> 2Br 2 +2H 2 0+0 2

Other reactions illustrative of this general principle are: 2C120 2C102

>60

°C

> 2C1 2 +0 2 ► Cl2+202

2C1207 Ξ225Ξ*22Ει^2α2+7θ2 > ~" 40 ° c > 2Br 2 +0 2 2Br20 2Mg(X03)2 ► 2MgO+50 2 +2X 2 (X = Cl, Br or I) > CsX+X2 (X = Br or I) CsX3 I2CI6 >2IC1+2C12

Laboratory Preparation and Purification Because of the ready availability of commercial chlorine, bromine and iodine, it is rarely necessary now to prepare these elements in the laboratory except for didactic reasons, though as recently as 1949 a method was described for the laboratory preparation of chlorine by electrolysis of chloride solutions50. Individual halogens can be satisfactorily separated from mixtures of the elements (and other volatile materials) by distillation or by gas-chromatographic51 methods. Chlorine, available in conveniently small cylinders, may be of high purity (99-9% or better), but certain commercial samples are liable to contain oxygen, nitrogen, carbon dioxide and monoxide, hydrogen chloride and moisture. A suggested52 method of purifi­ cation involves passing the chlorine first through potassium permanganate solution (to remove hydrogen chloride), then through sulphuric acid and over phosphorus pentoxide (to remove moisture), and finally fractionating the material in a vacuum-system, using either low-temperature baths [e.g. acetone-solid C 0 2 (-80°C), ethyl bromide (-119°C) and liquid nitrogen (—196°C)] or a low-temperature still. The reaction of chlorine both with mercury and with many kinds of tap-grease must be recognized in all manipulation of the purified material. Bromine of high quality is readily obtainable. The likely impurities, with the maximum limits imposed for AnalaR grade«, are as follows: Cl (0-025%), I (0-0001%), sulphate 50

W. J. Kramers and L. A, Moignard, Trans. Faraday Soc. 45 (1949) 903. si J. Janäk, M. Nedorost and V. Bubenikovä, Chem. Listy, 51 (1957) 890; J. F. Ellis and G. Iveson, Gas Chromatography (ed. D. H. Desty), p. 300. Butterworths, London (1958); J. F. Ellis, C. W. Forrest and P. L. Allen, Anal Chim. Ada, 22 (1960) 27. 52 R. E. Dodd and P. L. Robinson, Experimental Inorganic Chemistry, Elsevier (1954). 53 AnalaR Standards for Laboratory Chemicals, 6th edn., AnalaR Standards Ltd., London (1967).

FORMATION OF THE ELEMENTARY HALOGENS

1131

(0-005%), arsenic (0-0001%) and non-volatile matter (0-005%). The liquid (b.p. 58-8°C) can be distilled in a glass fractionating apparatus at atmospheric pressure or in vacuo*2*52; like chlorine, bromine attacks mercury, rubber and most tap-greases, which should therefore be excluded from the purification stages. Individual impurities can be removed by the following methods54: Impurity Chlorine Iodine Water Organic matter

Treatment Distillation from, or reaction with, a metal bromide; fractional condensation, low temperature-low pressure sublimation, or Chromatographie methods. Extraction with water or conversion to halide followed by treatment with perman­ ganate or bromine. Concentrated sulphuric acid or phosphorus pentoxide. (i) Saturation of a 5-30 % aqueous bromide solution with crude bromine, followed by distillation, condensation and drying of the bromine. (ii) Passage of bromine with oxygen at 1000°C through a furnace packed with frag­ mented quartz or high-silica (Vycor) glass, followed by removal of any chlorine and drying.

Very pure bromine for atomic weight determinations has been prepared55 by a lengthy process involving, inter alia, reduction of the bromine to bromide; oxidation of the iodide by permanganate; oxidation of the bromide to bromine with halogen-free potassium dichromate and sulphuric acid; distillation, drying over fused calcium oxide and calcium bromide, and finally vacuum-distillation of the bromine. Iodine of high purity can also be bought. The AnalaR grade has the following maximum limits of impurity«: Cl~ and Br~ (0-005%), C N " (0-005%), sulphate (0-01%) and non­ volatile matter (0-01 %). Further purification is effected by sublimation52, after first grinding up the iodine with a small amount of potassium iodide to reduce any free chlorine and bro­ mine and impurities such as IC1, IBr and ICN. Purification has also been accomplished32»33 by forming an insoluble heavy-metal iodide (Agl or Cul), and either (i) reducing the iodide with hydrogen to the metal and hydrogen iodide, and then oxidizing the hydrogen iodide with nitrite, or (ii) directly oxidizing the metal iodide to the oxide and iodine. Grease and mercury should, ideally, be excluded from the sublimation and other apparatus employed for handling purified iodine. Commercial Production of Chlorine, Bromine and Iodine Chlorine^^A6-4% The first patent connected with any industrial use of chlorine, dated 1799, was for its use in bleaching. During the last half of the nineteenth century the demand for chlorine for bleaching grew at a rate that brought about the invention and development of the Weldon and Deacon processes for chlorine production, in both of which hydrogen chloride (mostly from the sulphuric acid treatment of common salt) was the raw material. Towards the end of the nineteenth century, development of electrical generating equipment converted the electrolytic method of chlorine production from a laboratory to a commercial method. At the present time, in practically all countries, more than 99% of the chlorine of commerce is 54

V. A. Stenger, Angew. Chem., Internat. Edn. 5 (1966) 280. 5 O. Hönigschmid and E. Zintl, Annalen, 433 (1923) 201.

5

1132

CHLORINE, BROMINE, IODINE AND ASTATINE: A. J. DOWNS AND C. J. ADAMS

produced electrolytically, and chlorine ranks with sulphuric acid, sodium hydroxide, sodium carbonate, ammonia, phosphoric acid and nitric acid as one of the most important tonnage chemicals now made46. Chlorine, soda ash (Na 2 C0 3 ) and caustic soda (NaOH) represent the primary products of chemical industries whose basic raw material is common s a l t 46-48,56.

In the electrolysis of sodium chloride brines, chlorine is produced at the anode, while hydrogen is normally produced at the cathode, together with a simultaneous increase in the local hydroxide ion concentration. The overall reaction may be represented as Cl- +H 2 0 -> OH" +£C12 + £H2 In order to keep the anodic and cathodic products entirely separate, there are two distinct methods of operation: (i) the use of a diaphragm between the two compartments, and (ii) the use of a mercury cathode, whereby sodium, rather than hydrogen, is initially discharged. In the "diaphragm" process32»33»46-48, brine is fed continuously to the electrolytic cell, flowing from the anode compartment through an asbestos diaphragm backed by an iron cathode. To minimize back-diffusion and migration, the flow rate is always such that only part of the sodium chloride is converted to the hydroxide, hydrogen and chlorine. The catholyte solution of alkaline brine is evaporated to obtain the sodium hydroxide, in the course of which sodium chloride precipitates, is separated, redissolved and returned to the cell. The assembly consists of an outer steel shell, either cylindrical or rectangular, supporting the cathode, which may take the form of a perforated plate or woven screen inside the shell. The actual cathode surfaces are generally lined with a layer of asbestos, either in the form of paper or of vacuum-deposited fibres. At the minimum practical distance, the anodes are supported with their faces parallel to the diaphragm. An inert, insulating cover, carrying a brine inlet and a chlorine outlet, closes the cell. The choice of materials for the electrodes is important. For economic reasons, not only have corrosion and the consequent need for replacement to be reduced to a minimum, but also the hydrogen and chlorine overvoltages must be kept as low as possible to conserve energy. As cathode material iron or mild steel appears to be in general use, while graphite, with its virtues of moderate price, low apparent density and high overvoltage with respect to oxygen, is usually favoured for the anodes. However, the relatively high overvoltage of graphite with respect to chlorine and its limited mechanical strength and resistance to corrosion, which cause the anodes to disintegrate at a significant rate, have encouraged active interest in alternative anode materials, such as platinized titanium. In the mercury-cathode type of alkali-chlorine cell 32 » 33 · 46-48 , there are two main parts, the "electrolyser" or brine cell and the decomposer or "denuder". The "electrolyser" is essentially a closed rectangular trough, long in comparison with its height and breadth. Continuously fed brine is electrolysed between anodes, most commonly of graphite (though, again, platinized titanium offers notable advantages), and a cathode consisting of a sheet of mercury which flows uniformly over the flat bottom of the cell; chlorine gas is formed at the anodes and sodium amalgam at the cathode. Parts of the trough not in contact with mercury are provided with a corrosion-resistant, protective coating such as hard rubber; the bottom of the cell is in some cases bare steel and in others hard rubber, granite or concrete surfaced with a synthetic coating. The anodes are usually horizontal graphite plates suspended on 56

E. L. Gramse and L. H. Diamond, Advances in Chemistry Series, 78 (1968) 1.

1133

FORMATION OF THE ELEMENTARY HALOGENS

rods which extend through the cover of the cell; the anodes, parallel to, and close to, the mercury-brine interface, are perforated or grooved to facilitate the release of chlorine, which is removed from the cell either via an outlet in the cover or via an enlarged anolyte overflow connection. The sodium amalgam flows continuously to the decomposer, which may be regarded as a secondary cell where the amalgam becomes the anode to a shortcircuited graphite (or iron) cathode in an electrolyte of sodium hydroxide solution. There are two general types of amalgam decomposer: (i) a horizontal steel trough mounted parallel to the electrolyser containing graphite grids, and (ii) a vertical steel cylinder packed with pieces of graphite. Purified water is fed to the decomposer, generally in counter-current to the sodium amalgam; hydrogen gas is formed and the electrolyte increases in sodium hydroxide content. A solution containing from 30 to 70% sodium hydroxide at high purity overflows from the decomposer (most cells are operated to give a 50% solution of sodium hydroxide). The denuded mercury is recycled to the electrolyser; typically the amalgam reaching the decomposer contains 0-2% sodium and is returned with less than 0-02% sodium. For the overall aqueous reactions: (i) Diaphragm cell:

Cl ~ + H 2 0

> OH " + JC12 + £H2 HaO/Hg

+

(ii) Mercury cathode cell: Na +C\~

>Na/Hg+iCl2

are

the standard free energies, AG°298, (0 +50-445 kcal mol - 1 and (ii) approximately _1 + 76kcalmol corresponding to standard, reversible cell-voltages at 25°C of (i) 2-19 V and (ii) 3-3 V; the cell-voltages appropriate to the cells under typical operating conditions (given in Table 3) are necessarily somewhat different from these values. The reasons for the much higher voltages in the working cells are as follows: (i) there are high overvoltages of chlorine at the anodes of both cells and of hydrogen at the cathode of the diaphragm TABLE 3. TYPICAL OPERATING DATA FOR MODERN DIAPHRAGM AND MERCURY CELLS* - 0

Cell capacity Cell capacity Anode current density Cathode current density Current efficiency Cell voltage Energy consumption* Energy efficiency Cell temperature Graphite consumption Average anode life Diaphragm8, (or mercuryb) consumption Salt consumption NaCl concentration of inlet electrolyte NaCl concentration of exit electrolyte

Unit

Diaphragm cell

Mercury cell

amp kg Cl2/day amp c m - 2 amp c m - 2

10,000-60,000 300-2000 -01 ~01 96-5 3-75-3-85 -3-0 -58 95 2-7-3-6 220-310 0-6a 1-8 27 14-15

25,000-300,000 1000-10,000 ~0-5 0-35-Ό-5 94-97 4-25-4-50 -3-5 -47 65-70 2-0-2-8 150-450 0-2* 1-7 26 23-24

%

volt kWh/kg Cl 2

%

°C kg/1000 kg Cl 2 days kg/1000 kg Cl 2 kg NaCl/kg Cl 2 % by weight % by weight

* Electrolysis only a Z. G. Deutsch, C. C. Brumbaugh and F. H. Rockwell, Kirk-Othmer's Encyclopedia of Chemical Tech­ nology, 2nd edn., Vol. 1, p. 668, Interscience (1963). b C. A. Hampel (ed.), The Encyclopedia of the Chemical Elements, Reinhold, New York (1968). c J. S. Sconce (ed.), Chlorine: its Manufacture, Properties and Uses, American Chemical Society Mono­ graph No. 154, Reinhold, New York (1962).

1134

CHLORINE, BROMINE, IODINE AND ASTATINE: A. J. DOWNS AND C. J. ADAMS

cell (sodium has but a small overvoltage at the cathode of the mercury cell); (ii) there are potential drops through the electrolyte gap, through the diaphragm of the diaphragm cell, through the anode and cathode leads, and through the various contacts. The high overvoltage of hydrogen, relative to sodium, at a mercury cathode is one of the major principles embodied in the mercury cell; some 99% of the current brings about discharge of sodium ions and only about 1 % liberates hydrogen. In practice, 100% current efficiency is not realized in either type of cell for the following reasons: (i) There may be current leakages through insulators. (ii) Secondary reactions occur at the anodes: (a) 40H ~ (b) 40H " + C (c) 6C10 " + 3H 2 0

-> 2H 2 0+ 0 2 +4e -> 2H 2 0 + C0 2 +4e -> 2C103 " + 4C1" + 6H + + f0 2 + 6e

at the cathode: (d)C10-+2H + +2e -^C1"+H 2 0 (e) C10 3 -+6H + +6e-^Cl-+3H 2 0 (f) H + 4-e -> £H2 (in the mercury cell)

and in the anolyte or catholyte compartments (g) C12 + 20H(h) 3C12 + 60H(i) 2Na(Hg)+Cl2

^Cl-+C10-+H20 ^C10 3 -+5C1-+3H 2 0 ->2NaCl + Hg

The presence of certain impurities is likely also to affect the current efficiency. (iii) There are always leakages and losses of product. The relative merits of the two types of cells are nicely balanced, as may be judged in part from the data of Table 3. In practice, the diaphragm cell accounts for more than 70% of chlorine production in the United States, whereas in Canada and Western Europe the mercury cell is favoured. The two most important factors governing the choice between the cells are probably the physical form in which the salt is available and the quality of the sodium hydroxide required. Thus, the mercury cell requires dry salt whereas diaphragm cells are normally operated on saturated brine. In general, therefore, diaphragm cells are more economical where the operation is sited near a natural source of deep-well, saturated brine. On the other hand, the mercury cell produces sodium hydroxide of superior quality, though the factor of mercury-pollution via the effluent streams, the dangers of which have only recently, and rather sensationally, come to be appreciated57, may ultimately weigh against this type of cell. The auxiliary facilities for plants incorporating diaphragm or mercury cells, indicated in the flow diagrams of Fig. 9, include the following stages: brine purification; direct-current electric power supply; cooling, drying and liquefaction of chlorine; salt recovery (diaphragm cell only); cooling, filtering and concentration of the alkaline (caustic) solution; and cooling and compression of hydrogen, when the latter is utilized. By far the most important co-product is sodium hydroxide, whereas the hydrogen is a relatively minor factor, and much of it is simply burnt for fuel. The chlorine gas emerging from the cells is cooled, either by scrubbing with water or by contact with cooled stoneware, titanium or glass pipes, to remove water vapour and brine spray. It is then dried by a counter-current flow of sulphuric acid. 57 L. J. Goldwater, Scientific American, 224 (5) (1971) 15.

Cl, Free Vent-

τ

Weak H 2 S0 4 '

(a)

. Liquid C l t Product

CHLORINE | LIQUEFACTION

,Cl 2 Gas Product

CHLORINE PURIFICATION

CHLORINE COMPRESSION

■iConc. C^COOLING ΓΗΤΟ. & DRYING

CHLORINE RECOVERY

Anhydrous NaOH Products

CAUSTIC FUSION & FLAKING

. 73% NaOH Product

CAUSTIC CONCEN­ TRATOR

^50% NaOH Product

CAUSTIC COOLING & FILTERING

50% NaOH

I CAUSTIC & SALTEVAP | RECOVERY

Cell Liquor

Hydrogen Chlorine (cell gas)

Treating Chemicals

Hydrogen Product

L

HYDROGEN COMPRESSION

HYDROGEN COOLING

Hg-Na lAmalg.

NaOH Product

50% NaOH Product

Anhydrous NaOH Products

u

CAUSTIC FUSION & FLAKING

fc73%

CAUSTIC CONCENTRATOR

CAUSTIC FILTERING

ZZ3

BRINE DECHLQRIΝΑΤΙΟΝ

DECOMPOSER

,Η,Ο

ELECTROLYSIS

I Feed Fee

Brine ΙΒΓΪΙ

Hydrogen

CHLORINE RECOVERY

„ Liquid Cl 2 Product

CHLORINE LIQUEFACTION

Product

fcCl2Gas

CHLORINE COMPRESSION

Cl, COOLING &fH^SÖ I "DRYING I "

Chlorine (cell gas)

Treating BRINE Chemicals PURIFICATION

BRINE IRESATURATIONl

| Hydrogen

HYDROGEN COMPRESSION

HYDROGEN COOLING

FIG. 9. Flow diagrams (a) of diaphragm cell operation in a chlorine-caustic soda plant and (b) of mercury cathode cell operation in a chlorine-caustic soda plant. [Reproduced with permission from Kirk-Othmer's Encyclopedia of Chemical Technology, 2nd edn., Vol. 1, pp. 699, 700, Interscience, John Wiley and Sons, Inc. (1963).]

Salt

J[

ELECTROLYSIS

Brine Feed

BRINE IRESATURATIONl

Purified f Brine

BRINE TREATMENT

Salt or Brine

1136

CHLORINE, BROMINE, IODINE AND ASTATINE: A. J. DOWNS AND C. J. ADAMS

The concentration of the remaining impurities is further reduced by compression and, ultimately, liquefaction of the chlorine. Very corrosive to all of the common metals, wet chlorine is frequently handled in piping systems constructed of chemical stoneware, glass or some kinds of halogenated plastics. Dry gaseous or liquid chlorine can be handled satis­ factorily in steel apparatus at temperatures up to about 120°C. TABLE 4 . MISCELLANEOUS PROCESSES FOR THE MANUFACTURE OF CHLORINE* _ C

Principle Electrolytic oxidation of Cl~

Details (i) Aqueous potassium chloride (ii) Fused magnesium chloride

Chemical methods of oxidizing

ci-

Comments Operations similar to those involving sodium chloride brines Method of producing magne­ sium ; now puts little if any chlorine on the market

(iii) Fused sodium chloride (iv) Hydrochloric acid

Method of producing sodium Depends economically on the availability of cheap by­ product HC1

(i) Reaction of salt and nitric acid (ii) Catalytic air oxidation of HC1

Still practised on a small scale Variations on the Deacon Process; depends on the ready availability of cheap HC1

a Supplement to Mellor's Comprehensive Treatise on Inorganic and Theoretical Chemistry, Supplement II, Part I, p. 272, Longmans, Green and Co., London (1956). b Gmelins Handbuch der Anorganischen Chemie, 8 Auflage, "Chlor", Teil A, p. 2 (1968). c Z. G. Deutsch, C. C. Brumbaugh and F. H. Rockwell, Kirk-Othmer's Encyclopedia of Chemical Tech­ nology, 2nd edn., Vol. 1, p. 668, Interscience (1963).

Other processes that have been exploited in recent years for the technical production of chlorine32»33'46 are set out in Table 4. Perhaps the most significant change has been the increasing production of hydrogen chloride from chlorination or dehydrochlorination processes (as in the manufacture of solvents and synthetic resins), which has revived the problem of regenerating chlorine from hydrogen chloride: electrolysis of hydrochloric acid, catalytic oxidation of hydrogen chloride or metallic chlorides using air or oxygen, and the formation and catalytic decomposition of chlorosulphonic acid, CISO3H, represent, up to the present time, the principal lines of technical development. Brominei2>M>4S-w

All methods of bromine production depend on the oxidation of the bromide ion which is found naturally only in relatively low concentrations. In the commercial isolation of bromine from natural brines, bitterns or ocean water, four essential steps may be recognised: (a) oxidation of bromide to bromine; (b) removal of bromine vapour from the solution; (c) condensation of the vapour or fixation in some chemical form; and (d) purification of the product. Chlorine is the only oxidizing agent employed commercially in step (a), though

FORMATION OF THE ELEMENTARY HALOGENS

1137

some use has been made of electrolytic methods in the past. Step (b) involves driving out the bromine vapour with a current of air or steam. Steam is suitable when the raw brine is relatively rich in bromine (0-1% or more), but air is more economical when the bromine is extracted from very dilute solutions, such as ocean water58. When steam is used, the vapour may be condensed directly; otherwise the bromine must be trapped in an alkaline (sodium carbonate) or reducing (sulphur dioxide) solution. In either case, purification is necessary to remove chlorine. The so-called "steaming-out" method has been widely employed since the early German development of a continuous process and its further improvement in 1906 by Kubierschky. The Kubierschky process, in one form or another, is operated in Germany, Israel and the United States using brines that are relatively rich in bromide (e.g. mother liquors obtained during the processing of salts from the Stassfurt deposits, or liquors produced by solar evaporation of Dead Sea water). The so-called "blowing-out" process, in which air instead of steam is used to drive out the bromine, was developed in the United States in 1889 by Dow. This represents by far the most important process for obtaining bromine from sea water. It is operated on a large scale by the Dow Chemical Company at Freeport, Texas; other ocean-water plants are found at Hayle (Cornwall), Amlwch (Wales) and Marseilles (France). More detailed accounts of these processes are given elsewhere45. In a typical "steaming-out" plant raw brine is preheated to about 90°C and then passes into the chlorinator tower, a tank lined with resistant material and packed with rings. Only a portion of the total chlorine that is to be used is introduced into the bottom of the chlor­ inator. The brine then passes into the steaming-out tower where it is caused to flow uniformly over a packing made usually of chemical stoneware or porcelain. Steam is injected at the bottom of the tower and the remainder of the chlorine is introduced via a separate inlet higher up the tower. The weak acidity of the brine at this stage minimizes the hydrolysis of bromine and increases the efficiency of the steaming-out process; the brine leaving the steaming-out tower is neutralized with lime and finally pumped out through a heat-exchanger. The halogen-laden vapour passes into a condenser and then into a gravity separator. Non-condensable gases bearing some chlorine and bromine are returned to the bottom of the chlorinator tower, while the water layer from the separator, saturated with chlorine and bromine, is returned to the steaming-out tower. Crude bromine collects at the bottom of the separator, whence it passes through a trap to the first of two distillation columns, in which the free halogens are separated from higher-boiling halogenated hydrocarbons. The chlorine from this stage is recycled to the steaming-out tower, while the bromine undergoes a second fractionation for final purification. More than 95% of the bromine is thus recovered from the brine. By contrast, the following operations are typical of the "blowing-out" process. Sea water is pumped to the top of a brick and concrete blowing-out tower; on the ascent, it is treated with sufficient sulphuric acid to bring the pH down to 3*5 and with a quantity of chlorine 15% in excess ofthat theoretically required. The pipe through which the water ascends is lined with rubber, as are also the injection pipes for acid and chlorine. At the top of the tower the liquid is distributed so that it descends through parallel chambers filled with wood packing. Here the free bromine is given up to a counter-current stream of air drawn in at the base of the tower. The mixture of air and halogen vapour is then caused to 58 W. F. Mcllhenny, Mater. Technol.:—Interamer. Approach, Interamer. Conference (1968) 119; Chem. Abs. 70 (1969) 40537g.

1138

CHLORINE, BROMINE, IODINE AND ASTATINE: A. J. DOWNS AND C. J. ADAMS

react with sulphur dioxide and water vapour: X2 +S0 2 +2H 2 0 -> 2HX+H2S04 (X = C\ or Br) The spray of mixed acids is collected and transferred to a steaming-out tower in which the bromine is liberated by the action of chlorine, and collected as in the steaming-out process. The sulphuric and hydrochloric acids formed are used for acidifying the sea water in the first stage of the process. The whole process is automatically controlled, the rate of addition of chlorine being governed by measurement of the oxidation potential of the liquid with a platinum electrode. A high initial temperature of the sea water clearly favours extraction of the bromine by diminishing its solubility in water, and this consideration accounts for the location of some of the plants. Before the general introduction of sulphur dioxide for the removal of bromine from the air stream, absorption of the halogen was usually carried out using sodium carbonate solution : 3C0 3 2 " + 3Br2 -> 5Br " + Br0 3 ~ + 3C0 2 3C0 3 2 " +fCl 2 +£Br 2 -> 5C1" + B r 0 3 " + 3 C 0 2

Acidification of the resulting solution of sodium bromide and bromate with sulphuric acid regenerates free bromine which can be steamed out and condensed in the manner already described. This alkaline absorption procedure, in one form or another, is still used in some countries. Some of the other processes which have been used in the extraction of elementary TABLE 5. MISCELLANEOUS PROCESSES FOR THE EXTRACTION OF BROMINE* -C

Principle

Method

Comments

Electrochemical oxidation of aqueous bromide

In general terms as for chloride with either a diaphragm (as in the Wünsche process) or special bipolar carbon electrodes (the Kossuth process)

Now generally uneconomic mainly because of the huge volumes of brine which have to be handled

Chemical oxidation of bromide

Use of manganese dioxide or chlorate under acidic conditions

Early processes for the production of bromine, now obsolete

Alternative methods for the fixa­ tion of bromine liberated by blowing-out methods

(i) Reaction with moist scrap iron, ammonia or ferrous bromide solutions (ii) Interaction of bromine and phenol to give tribromophenol, followed by fusion of the latter with alkali (iii) Interaction of bromine and aniline to give tribromoaniline (iv) Solvent extraction; absorp­ tion by silica gel or activated charcoal; ion exchange

No longer used to produce elementary bromine; may be used to produce bromides A process developed in Russia

a

Yields relatively poor; tried as a venture by the Ethyl Corporation Methods suggested, e.g. in the patent literature

Z. E. Jolles (ed.), Bromine and its Compounds, Benn, London (1966). V. A. Stenger and G. J. Atchison, Kirk-Othmer's Encyclopedia of Chemical Technology, 2nd edn., Vol. 3, p. 750, Interscience (1964). c C. A. Hampel (ed.), The Encyclopedia of the Chemical Elements, Reinhold, New York (1968). b

FORMATION OF THE ELEMENTARY HALOGENS

1139

bromine are listed in Table 5. Some attention has been given to new ways of isolating bromine, including, for example, the use of ion exchange and solvent extraction59, but these methods have not come into commercial use. It may be possible to avoid the need for chlorine in the regeneration of bromine from the aqueous solution that results from the interaction of the halogen vapours with sulphur dioxide; a new process entails the functioning of a trace of butyl nitrite or an alkali-metal nitrite in an acidic medium as a catalyst for oxidation by oxygen or air at 0° to 200°C and pressures of 1-4 atm6<>. When the blowing-out process is applied to brine with the object of manufacturing not elementary bromine, but alkali or alkaline earth bromides and bromates, several variations of procedure are advantageous45 _47 . In particular, the air-halogen mixture is usually caused to react with the appropriate alkali carbonate or alkaline-earth hydroxide to give a mixture of bromide and bromate: 3Br2 + 3H 2 0 -> 5Br " + Br0 3 ~ + 6H +

Subsequent treatment depends upon the compound desired. Thus, bromates of sodium and potassium are less soluble than the bromides and may be crystallized from the mixtures after partial evaporation and cooling. The mother liquors rich in bromide may be heated with scrap iron to reduce the bromate and the bromides crystallized, after removal of the iron oxide precipitate. Other methods of removing bromate include (a) heat treatment, preferably in the presence of charcoal, (b) precipitation of sparingly soluble barium bromate, (c) reduction with hydrogen (or barium) sulphide, and (d) reaction of the bromine with ammonia, whereby bromate formation is suppressed: 8NH 3 + 3Br2 -> 6NH 4 Br+N 2

Equipment coming into contact with wet bromine is generally made of ceramio materials45 - 4 7 . Glass and tile piping are employed most extensively, though some use has been made of tantalum metal, particularly for condensers. The dry halogen may be handled in lead (or lead-lined) or Monel equipment. Brick, granite or concrete construction is favoured for large tanks or other vessels. /ο<#Λί?32,33,46,47

For the production of iodine from natural brines, one of several possible processes may be used. The first step in any such process is the clarification of the brine to remove oil and other suspended material. In one process, a silver nitrate solution is added in sufficient quantity to precipitate only silver iodide, which is filtered off and treated with clean steel scrap to form metallic silver and a solution of ferrous iodide. The silver is redissolved in nitric acid for use in another cycle and the solution is treated with chlorine or some other oxidizing agent to liberate the iodine. In another process, the step after clarification is treatment of the brine with chlorine gas to liberate free iodine. The solution is then passed over bales of copper wire with which the iodine reacts to form insoluble cuprous iodide. At intervals the bales are agitated with water to separate the adhering iodide and returned to the precipitators for further use. The cuprous iodide suspended in the water is filtered, dried and stored or transported as such. 59 R. F. Hein, U.S. Pat. 3,037,845 (1962); 3,174,828 (1965); F. J. Gradishar and R. F. Hein, U.S. Pat. 3,098,716 (1963); L. C. Schoenbeck, U.S. Pat. 3,075,830 (1963); 3,101,250 (1963); E. I. du Pont de Nemours; M. Israel, A. Baniel and O. Schachter, Israel Pat. 15,408 (1962). 60 W. A. Harding and S. G. Hindin, U.S. Pat. 3,179,498 (1965), Air Products and Chemicals, Inc., and Northern Natural Gas Co.

1140

CHLORINE, BROMINE, IODINE AND ASTATINE: A. J. DOWNS AND C. J. ADAMS

The largest part of the iodine produced in the United States is derived as a by-product of the Michigan natural brines, from which the iodine is recovered by a process resembling that for the recovery of bromine from sea water. The brine, containing typically 30-40 ppm iodine, is acidified with sulphuric acid and treated with a slight excess of chlorine to liberate the iodine. The liquor is then pumped to a denuding tower in which it is stripped of its iodine by a counter-current of air. The air passes to a second tower where the iodine is absorbed by a solution of hydriodic and sulphuric acids into which sulphur dioxide is passed continuously to reduce the iodine to iodide. This solution, when sufficiently concentrated, is again chlorinated; the iodine precipitated is filtered off, melted under concentrated sulphuric acid and cast into pigs. Brines may also be treated with a mixture of sulphuric acid and sodium nitrite: 21 + 2 N 0 2 + 4 H + - ^ I 2 + 2 N O + 2 H 2 0

The iodine liberated is conveniently extracted, for example, with petroleum or activated charcoal, and converted back to iodide by treatment either with sulphite or with alkali. Oxidation of the iodide concentrate to elementary iodine may be effected with sodium dichromate and sulphuric acid; the use of other oxidizing agents, e.g. bleaching powder, has also been investigated. Electrolytic methods for the extraction of iodine have generally proved to be commercially impractical because of the very low concentrations of iodide present in natural brines. Recently developed ion-exchange techniques have been reported to give a new, more economical extraction process for producing iodine of exceptionally high purity47. The process used by the Chilean nitrate industry differs from all the others since the natural source of iodine is here iodate. Sodium iodate is extracted from the caliche, being -allowed to accumulate in the mother liquors from crystallization of sodium nitrate until a suitable concentration (about 6 g 1_1) has been attained. Part is then drawn off and treated with the exact quantity of sodium bisulphite required to reduce all of the iodate to iodide: 2 I 0 3 " + 6HSO3 - -> 21" + 6SO42 " + 6 H +

The resulting acidic mixture is then caused to react with just sufficient fresh mother liquor to liberate all the iodine in accordance with the reaction, 5I-+IO3 -+6H + ^3I 2 +3H 2 0 The precipitated iodine is filtered off in bag filters and the iodine-free solution returned to the nitrate-leaching cycle after neutralization of any excess acid with sodium carbonate. The iodine cake is washed with water, pressed to reduce the moisture content, broken up and sublimed in concrete-lined iron retorts connected to condensers made of glazed tile. The product is crushed and packed in polythene-lined fibre drums. In France, Russia and Japan where alginic products are prepared from seaweed, small quantities of iodine can be obtained at low cost as a by-product. The stages in the prepara­ tion are: drying and burning the seaweed, leaching of the ash, release of the iodine by chemical reaction (e.g. with Mn0 2 ), and purification. 2.4. U S E S OF C H L O R I N E , B R O M I N E A N D IODINE32.45-48

Approximate data on the annual production of chlorine, bromine and iodine, together with the principal commercial uses of the elements, are presented in Table 6. Originally the

ATOMIC CHLORINE, BROMINE AND IODINE

1141

primary commercial outlet for chlorine, the bleaching of textiles now accounts for less than 1% of the total output of chlorine. In fact, chlorine is too destructive for bleaching wool, silk and other products of animal origin, and in recent years hydrogen peroxide has become economically available for these and other purposes. The rapid growth in chlorine usage in recent years (implied by the thirtyfold increase in production in the period 1930-50 and the further tenfold increase in the period 1950-7032»46,47) i s to be correlated with the growth of the organic chemical industry, and especially with the production of compounds derived from petroleum. Thus, about 65% of the chlorine produced in the United States in 1963 was used for the manufacture of organic chemicals important, for example, as intermediates, solvents, refrigerants, plastics and plasticizers, insecticides and dyes46. To complete the pattern of chlorine usage for the United States in 1963, 9% went into the production of inorganic chemicals, e.g. hydrochloric acid, bromine, sodium hypochlorite and various metal chlorides, 17% into the treatment of wood pulp in the production of paper and rayon, 4% into the disinfection of water and the treatment of sewage, and 5% into a variety of minor applications. Only about 3% of commercial bromine now finds its way into inorganic bromides and bromates, in complete contrast to the early economic history of bromine, which was closely linked to these derivatives45»46. By far the most important use of bromine (on a tonnage basis) is in the manufacture of ethylene dibromide for motor fuels; the dibromide serves to convert the lead from the tetraalkyllead antiknock agent to lead(II) bromide, which is volatile at cylinder combustion temperatures and is swept out with the exhaust. The great bulk of commercial bromine is otherwise converted into bromine compounds, as there are relatively few direct applications of liquid bromine outside the laboratory. Organic bromine derivatives find uses, inter alia, as industrial organic intermediates, dyestuffs, pharmaceutical chemicals, for fumigation and pest control, as high-density liquids in processes deploying gravity separation, and as fire-extinguishing and fire-proofing agents. Of inorganic bromine compounds, alkali-metal bromides have applications in medicine and (as sources of hypobromite) in bleaching and the desizing of cotton, while silver bromide is a primary agent in photography, and bromates are employed as oxidants for specific purposes. Unlike chlorine and bromine, iodine has no major commercial outlet, though over half of the iodine produced is consumed industrially, and important quantities also find uses in nutrition, sanitation and medicine. Some of the more significant uses of iodine and its compounds depend upon (a) their function as catalysts, (b) the disinfectant properties of the element and the nutritional properties of iodides, (c) the photosensitive behaviour of silver iodide, a constituent of rapid negative emulsions in photography, and (d) the colour and other properties of iodine-containing dyestuffs.

2.5. A T O M I C C H L O R I N E , B R O M I N E A N D IODINE32.44.6i

Formation In conventional circumstances, halogen atoms are thermodynamically unstable and short-lived with respect to the corresponding diatomic molecules. Nevertheless, there has been a long and sometimes chequered history concerning the formation and characterization of chlorine, bromine and iodine atoms. By making vapour-density measurements at high 6i W. A. Waters, The Chemistry of Free Radicals, 2nd edn., Clarendon Press, Oxford (1948).

(i) The production of organic chemicals, the main uses of which are as (a) solvents, (b) refrigerants and dispersion propellants, (c) plastics, resins, elastomers and plasticizers, (d) intermediates in the production of deter­ gents and other materials, (e) insecticides, (f) dyes, (g) automobile antifreeze and anti­ knock compounds. Some important organic compounds so pro­ duced are: Aliphatic: Cx: CH3C1, CH 2 C1 2 , CHCI3, CCI4; C 2 : ethylene oxide, ethylene dichloride; di-, tri- and tetra-chloroethylene, vinyl chloride; ethyl chloride, tetraethyllead; monochloroacetic acid, chloral; di(chloroethyl) ether; C3: glycerol, acrylonitrile; propylene glycol: C4 and larger fragments'. butadiene, chloroprene; chloro-derivatives of C5 and larger fragments. Aromatic: Chlorobenzene, phenol, aniline, dodecylbenzenes, diphenyl, diphenyl ether; chlorotoluenes, 0- and /7-dichlorobenzene, DDT; 2,4-dichlorophenoxyacetic acid, hexachlorobenzene, pentachlorophenol. (ii) The production of inorganic materials: HC1, metal chlorides, e.g. aluminium trichloride and titanium tetrachloride; sodium hypochlorite and bleaching powder; Br 2 and I 2 . (iii) Treatment of wood pulp (to remove lignin and other undesirable constituents); bleach­ ing of pulp and paper; textile bleaching.

Uses

1

1 -5-2-0 x 105 metric tons per annum

1 -5-2-0 x 107 metric tons per annum

Estimated world production

5-9 x 103 metric tons per annum

Iodine

(i) The manufacture of ethylene dibromide for (i) The element. Used (a) as a catalyst, e.g. in the conversion of resins, tall oil and other antiknock motor fuels. wood products to more stable forms, and of (ii) The production of organic derivatives such as amorphous to black selenium; (b) as a methyl bromide, ethylene dibromide and selective absorption agent, e.g. of olefinbromochloropropanes for insect control, as parafRn hydrocarbon mixtures; (c) to give space and soil fumigants and fungicides. charge-transfer complexes with aromatic (iii) Formation of industrial organic intermedi­ hydrocarbons which are effective as lubri­ ates, dyestuffs, medicinal chemicals and cants; (d) for the control of bacterial growth solvents, e.g. bromoanthraquinones, bromoin cutting-oil emulsions and for disinfection phthaleins, CF3CHBrCl ("halothane") and of water and food; (e) "tincture of iodine" mercurochrome. and "iodophors"; (f) in quartz-iodine lamps. (iv) As a disinfectant and bleaching agent. (v) Production of organic liquids of high density (ii) Inorganic iodides: (a) Agl, used in photog­ raphy and for the seeding of clouds to for gauge fluids and gravity separation pro­ induce rainfall; (b) certain iodides used as cesses, e.g. CBr4, CHBr 3 , CH 2 Br 2 , catalysts, e.g. Nil 2 in the addition of CO to Br 2 CHCHBr 2 , BrCH 2 CH 2 Br and C 6 H 5 Br. organic compounds and T1I4 in the pro­ (vi) Manufacture of fire-extinguishing agents, duction of stereospecific polymers; (c) in the e.g. CH2ClBr and CF 2 Br 2 , and of flame form of iodized salt and protein, which are retardants for incorporation in polymers essential to the health and well-being of such as polystyrene, e.g. man and the higher animals; organic and CH2Br[CHBr]2CH2Br, BrCl 2 CCCl 2 Br, inorganic derivatives of radioactive131I have pentabromochlorocyclohexane and tetrabeen used in the treatment of certain bromoisopropylidenebisphenol. thyroid and heart conditions; (d) addition of (vii) Production of inorganic bromides and broCdl 2 or Pbl 2 to electric generator and mates: AgBr, central to photography; alkali motor brushes to improve their efficiency bromides used as mild sedatives; with aqueous and life; (e) the production of high-purity hypochlorite bromide gives a solution con­ metals or metalloids, e.g. Ti, Zr, Hf and Si, taining hypobromite useful for bleaching by the van Arkel method. and for desizing cotton; hydrogen bromide, used as a catalyst, e.g. in alkylation reac­ (iii) Inorganic iodates: NaIC>3, improves the tions; bromates, used as oxidants for bread-making qualities of certain flours.

Bromine

Chlorine

TABLE 6. PRODUCTION AND USES OF CHLORINE, BROMINE AND IODINE* ~ Θ

improving the baking quality of wheat flour, (iv) Organic iodides: (a) Dyes, e.g. diiodoin certain hair-wave preparations, and fluorescein, rose bengal and erythrosin, used during the malting process in the brewing in photographic processes and in food industry. (All of these probably relate to the colouring agents; (b) high-density liquids oxidation of -SH to -S-S- groups.) for gravity separation (mostly in the labora­ tory), e.g. CH2I2; (c) X-ray contrast media, e.g. sodium 3-acetylamine-2,4,6-triiodobenzoate; (d) germicidal agents, e.g. CHI3; (e) intermediates, e.g. CH3I and C2H5I.

Supplement to Mellofs Comprehensive Treatise on Inorganic and Theoretical Chemistry, Supplement II, Part I, Longmans, Green and Co., London (1956). Z. E. Jolles (ed.), Bromine and its Compounds, Benn, London (1966). Z. G. Deutsch, C. C. Brumbaugh and F. H. Rockwell, Kirk-Othmer 's Encyclopedia of Chemical Technology, 2nd edn., Vol. 1, p. 668. Interscience (1963); V. A. Stenger and G. J. Atchison, ibid., Vol. 3, p. 750 (1964); A. W. Hart, M. G. Gergel and J. Clarke, ibid., Vol. 11, p. 847 (1966). d C. A. Hampel (ed.), The Encyclopedia of the Chemical Elements, Reinhold, New York (1968). β E. L. Gramse and L. H. Diamond, Literature of Chemical Technology, p. 1 (Advances in Chemistry Series Vol. 78), American Chemical Society, Washing, ton D.C. (1968).

a b c

(iv) Sanitation and disinfection: Treatment of water supplies and sewage; sterilization of process liquors in sugar manufacture and of sea-water cooling circuits. (v) Extraction and refining of metals: Orebeneficiation, fluxing; involved in methods for the extraction of such metals as Ti, Zr, Cu, Pb, Zn, Ni, Au, W and V, and for the refining of metals, e.g. AI, Mg and Pb.

1144

CHLORINE, BROMINE, IODINE AND ASTATINE: A. J. DOWNS AND C. J. ADAMS

temperatures, Victor Meyer62 showed that the I2 molecule dissociates reversibly thus: I2 ^ 21. In 1887 J. J. Thomson reported63 that bromine and iodine vapours could be dissociated by electrical sparks, though subsequent experiments64 failed to reproduce the changes of vapour-density which he described. Since 1916 chemists have correctly inferred that some atomic chlorine is produced when chlorine gas is exposed to ultraviolet light, a hypothesis which has subsequently proved a consistent prerequisite to the mechanistic interpretation of chlorination reactions in the gas phase. It was shown in 1928 that traces of atomic chlorine are produced by passing sodium vapour into chlorine gas61 and that the combination of hydrogen and chlorine gases can thereby be initiated in the absence of light. Appreciable quantities of atomic chlorine and bromine were first obtained in 1933 by the action of an electrical discharge on the appropriate molecular gas65»66. The following methods are now generally acknowledged to be effective in the production of halogen atoms. Thermal Dissociation The latest thermodynamic data67 indicate (see, for example, Table 7) that at temperatures TABLE 7. THERMAL DISSOCIATION OF HALOGEN MOLECULES»

Chlorine

Temperature (°K)

logtf p

298 500 1000 1500 2000 3000

-36-800 -19-626 -6-796 -2-452 -0-256 1-964

Degree of logtf p dissoc. (%)* 40x10-17 1-5x10-8 40x10-2 5-8 51-8 99

Iodine

Bromine Degree of dissoc. (%)*

-28-880 3-6x10-13 -14-660 4-7x10-6 -4-490 0-57 25-6 -1054 85 0-692 2-480 100

log#p

Degree of dissoc. (%)*

-24-628 -10-510 -2-524 0172 1-536 2-928

4-9x10-11 5-6x10-4 40 68-5 98 100

* At a pressure of 1 atm. JANAF Thermochemical Tables, The Dow Chemical Company, Midland, Michigan (1960-8). a

in the range 60O-900°K and pressures near 1 atm the gas-phase dissociation X2^2X

begins to become appreciable (degree of dissociation > 0-01%): at 1 atm 1% dissociation is achieved at about 1250°, 1050° and 850°K for Cl2, Br2 and I2, respectively. Rapid, homogeneous heating can be achieved by the use of shock waves68, following which the rates of dissociation and recombination of the halogens have been investigated by 62 V. Meyer, Ber. 13 (1880) 394. 63 J. J. Thomson, Proc. Roy. Soc. 42 (1887) 343. 64 E . P. Perman, Proc. Roy. Soc. 48 (1890) 4 5 ; W. Kropp, Z. Elektrochem. 21 (1915) 356. 65 G . - M . Schwab and H . Friess, Naturwiss. 21 (1933) 2 2 2 ; Z . Elektrochem 39 (1933) 586. 66 W. H . Rodebush and W. C. Klingelhoefer, jun., / . Amer. Chem. Soc. 55 (1933) 130; G . - M . Schwab Z . physik. Chem. B27 (1934) 4 5 2 . 67 JANAF Thermochemical Tables, The D o w Chemical Company, Midland, Michigan (1960-1968). Reactions 68 E . F . G r e e n e , / . Amer. Chem. Soc. 76 (1954) 2127; E . F . Greene and J. P. Toennies, Chemical in Shock Waves, Edward Arnold, London (1964).

ATOMIC CHLORINE, BROMINE AND IODINE

1145

measuring the optical absorption as a function of time. Thus, for mixtures of 5% Cl2 and 95% Ar at temperatures between 1600° and 2600°K, the rate of dissociation is given69 by log kj) (mol - 1 1 sec -1 ) = 10-66-9930/Γ, corresponding to an apparent activation energy, EA, of 45 kcal mol - 1 ; for comparable Br2/Ar mixtures™ at temperatures between 1500° and 1900°K, log kD = 10-709-8240/Γ, whence EA = 38 kcal mol"*. Dissociation by Discharge The dissociation of halogen molecules into atoms can also be effected in the gas phase by the action of a discharge. Arrangements have been described for the production of atomic chlorine or bromine in which an electrical discharge is maintained either externally65 or between water-cooled iron electrodes situated inside the reaction vessel66. The conditions normally favoured include ambient temperatures and gas pressures < 1 mm Hg, leading to reported degrees of atomization of 10-40%. Relatively high concentrations of chlorine or bromine atoms have also been generated in a flow system by means of microwave71 or radiofrequency72 discharges. Curiously, attempts to obtain iodine atoms in this way have been unsuccessful71; this difficulty is presumably a function of the hot region which develops immediately after the discharge, because it has been found that iodine atoms can be main­ tained if they are created outside the discharge. At ~ 0-1 mm pressure, the mean lives of chlorine and bromine atoms in a glass tube have been estimated65»66 to be about 3 x 10 ~3 and 1-8 x 10 ~3 sec, respectively. Reversion of the atoms to the molecular state is a relatively slow and inefficient process because a termolecular collision is necessary to dissipate the energy of combination: Χ·+Χ·+Μ->Χ 2 +Μ* and the rate of recombination is largely determined by the effectiveness of the "third body" M as an energy sink, a role which has been investigated for many materials. In effect, recombination normally takes place on the surfaces of the reaction vessel. On a dry quartz surface it has been estimated73 that the activation energy of recombination of chlorine atoms is ~ 1 kcal mol - 1 , and that, on average, about 1 in 12 collisions with the surface leads to union of the atoms (cf. the report that only 1 in 105-109 collisions between pairs of atoms in the gas phase leads to combination). Most metal surfaces are efficient "third bodies" for the recombination of halogen atoms, and a thin silver mirror has been used as a sensitive test for halogen atoms, which, in contrast with the corresponding molecules, cause immediate discoloration66»71. However, glass surfaces can be poisoned by certain materials, e.g. H 2 S0 4 , H3PO4, H 3 As0 4 , H3BO3 and HC104, which inhibit the rapid recombination of halogen atoms and so permit the achievement of relatively high concentrations of these atoms in a suitable flow system71. Optical Dissociation The absorption spectra of gaseous chlorine, bromine and iodine exhibit in the visible region a series of bands sharply degraded to longer wavelengths with convergence limits at 4795, 5108 and 4991 Ä, respectively74, beyond which the absorption becomes continuous. 69 M. van Thiel, D . J. Seery and D . Britton, / . Phys. Chem. 69 (1965) 834; see also R. A. Carabetta and H. B. Palmer, / . Chem. Phys. 46 (1967) 1333. 70 C. D . Johnson and D . Britton, / . Chem. Phys. 38 (1963) 1455. 71 E. A. Ogryzlo, Canad. J. Chem. 39 (1961) 2556. 72 V. Beltran-Lopez and H. G. Robinson, Phys. Rev. 123 (1961) 161. 73 G.-M. Schwab, Z. physik. Chem. A178 (1936) 123. 74 A. G. Gaydon, Dissociation Energies, 3rd edn., Chapman and Hall, London (1968).

1146

CHLORINE, BROMINE, IODINE AND ASTATINE: A. J. DOWNS AND C. J. ADAMS

The convergence limit corresponds in each case to the threshold energy of the dissociation Χ2(1Σ,+)^Χ(2Ρ3/2)+Χ(2ΡΙ/2)

With both bromine and iodine it has been possible to calculate the chemical composition of the photo-stationary state of equilibrium which is reached, under constant illumination, between the dissociation and recombination processes75: ->Χ·+Χ· (a)X2+/n> (b)X-+X-+M->X 2 +M* The energy transmitted to the "third body" M in the recombination reaction is in part converted into heat, with the result that when a photochemically dissociable gas, such as chlorine, bromine or iodine, is first illuminated, there is an increase of pressure, a phenom­ enon known as the Budde effect, arising from the gain in the total kinetic energy of the gas particles32. Decisive chemical evidence of the formation of free atoms upon irradiation of the halo­ gens is afforded by a wealth of facts. Thus, the photochemical reactions of chlorine or bromine with hydrogen and of iodine with olefins76 are satisfactorily explicable only in terms of photodissociation of the halogen molecules. By the methods of flash photolysis77, halogen atoms at pressures up to several cm Hg have been produced and detected by their absorption spectra; in these circumstances the half-life of chlorine atoms has been estimated to be ~ 3 x 10 ~2 sec. Flash photolysis has also provided a convenient and elegant method of studying the recombination reactions, notably of iodine atoms77. As a rule, the weaker heteronuclear bonds in which a halogen is engaged are susceptible to photolytic dissociation, leading commonly to the formation of free halogen atoms, e.g. hv

RI

* R -+I-61 hv

R · O · Cl

> RO + Cl ·

(R = organic group)™

Neutron Irradiation Neutron irradiation of halogen-containing species has been shown commonly to proceed with homolytic cleavage of chemical bonds to produce radioactive halogen atoms, which normally react with their chemical environment (see pp. 1169-72)32,45,79-81. Formation as Intermediates in Chemical Reactions The formation of free halogen atoms is believed to be an essential stage in numerous halogenation, dissociation and other reactions, on the evidence usually of detailed studies of the kinetics and conditions of the reactions. Examples are: H + x

2

^HX+X·

R · + Cl 2 -> R-Cl + Cl ·

(R = organic group)

Cl · + XY -> XC1+Y ·

(XY = Br 2 , ClBr or C1I)82

75 E . R a b i n o w i t c h a n d H . L . L e h m a n n , Trans. Faraday Soc. 3 1 (1935) 6 8 9 ; E . R a b i n o w i t c h a n d W . C . W o o d , ibid. 3 2 (1936) 5 4 7 · 76 G . S. F o r b e s a n d A . F . N e l s o n , / . Amer. Chem. Soc. 5 9 (1937) 693. 77 R . G . W . N o r r i s h , Angew. Chem. 6 4 (1952) 4 2 1 ; G . P o r t e r , Proc. Roy. Soc. 200A ( 1 9 5 0 ) 2 8 4 ; R . G . W . N o r r i s h a n d B . A . T h r u s h , Quart. Rev. Chem. Soc. 10 (1956) 149. 78 W . A . P r y o r , Free Radicals, M c G r a w - H i l l , N e w Y o r k (1966). 1 (1959) 267. 79 G . H a r b o t t l e a n d N . Sutin, Adv. Inorg. Chem. Radiochem. so R . Wolfgang, Ann. Rev. Phys. Chem. 16 (1965) 15. 81 R. Wolfgang, Progress in Reaction Kinetics, Vol. 3 (ed. G. Porter), p. 97, Pergamon (1965). 82 M . I . Christie, R . S. R o y a n d B . A . T h r u s h , Trans. Faraday Soc. 5 5 (1959) 1139, 1149.

ATOMIC CHLORINE, BROMINE AND IODINE Br· + X Y -> X B r + Y ·

1147

(XY = Cl 2 or BrCl)82

ArN 2 Cl

-> Ar · + N 2 + C l · (Ar = aryl group) Δ ci. CO + C1· < COC1C *OCl2+Cl·

NO + Cl 2 - * NOC1+Cl ·

Ι

(ref. 83)

Δ

NO +C1 Cl · + ClO « ^ - C l 2 - ^ ClO + Cl · I\ 03 + 02

o

(ref. 84)

N

| Cl+20 2 ci· + o 2 H · + HI -> H 2 +I*(2P 1 / 2 )

(ref. 85)

CuCl»-(»-2) ^ CuCl«_i-( w -2) + Cl· aq soln HO+H30++Cl*2H20+Cl·

(ref. 86) (ref. 87)

Solvated halogen atoms are also generated, it is believed, by the action of ionizing radiation (neutrons, α-particles, electrons, X- or y-rays)88 or of ultraviolet light89 on aqueous solutions of halide ions, a scheme such as hv X~aq^X~aq

- * [X* aq + e~aq]

being proposed to account for the photochemistry of chloride, bromide and iodide ions in aqueous solution90. Detection and Characterization Apart from the qualitative detection of halogen atoms based for example on the dis­ coloration of a metal mirror?! or on the observed characteristics of a particular reaction, the concentration of the atomic species, and hence the rate of their recombination and other reactions, may be measured by various methods. Notable among these are spectrophotometry, most commonly in the visible region, mass spectroscopy, thermal conductivity measurements, the use of a thermocouple detector (which depends on the heating effect of the recombination of the atoms on a metal wire)91 and chemical titration procedures involving some rapid reaction such as Cl-+NOCl-*NO+Cl27i The characterization of atoms of the individual halogens turns for the most part on their spectroscopic properties. Thus, the esr spectra of atoms produced in the gas phase have been observed for chlorine, bromine and iodine in their ground (2P3/2) states92, as well as for 83 P. G . Ashmore, Trans. Faraday Soc. 4 9 (1953) 2 5 1 ; P. G . Ashmore a n d J. Chanmugam, ibid. 2 5 4 . 84 H . Niki and B . W e i n s t o c k , / . Chem. Phys. 4 7 (1967) 3 2 4 9 ; P. H u h n , F . Tudos and Z. G . Szabo, M.T.A. Kam. Oszt. Közl. 5 (1954) 4 0 9 . 85 P. Cadman and J. C . Polanyi, / . Phys. Chem. 7 2 (1968) 3715. 86 J. K . Kochi, / . Amer. Chem. Soc. 8 4 (1962) 2121. 87 H . Taube and W . C . Bray, / . Amer. Chem. Soc. 6 2 (1940) 3357. 88 A . O. Allen, C . J. Hochanadel, J. A . Ghormley and T. W . D a v i s , / . Phys. Chem. 5 6 (1952) 575; M . Cottin, J. chim. Phys. 5 3 (1956) 903. 89 J. Jortner, M . Ottolenghi and G . Stein, / . Phys. Chem. 6 8 (1964) 247. 90 M . F . F o x , Quart. Rev. Chem. Soc. 2 4 (1970) 565. 91 J. W. Linnett and M. H. Booth, Nature, 199 (1963) 1181. 92 N . Vanderkooi, jun., and J. S. MacKenzie, Advances in Chemistry Series, 3 6 (1962) 9 8 ; S. Aditya and J. E . Willard, / . Chem. Phys. 4 4 (1966) 8 3 3 ; E . Wassermann, W. E . Falconer and W. A . Yager, Ber. Bunsenges. Phys. Chem. 7 2 (1968) 2 4 8 .

1148

CHLORINE, BROMINE, IODINE AND ASTATINE! A. J. DOWNS AND C. J. ADAMS

chlorine in its excited (2Λ/2) state93, though attempts to detect the atoms trapped in solid, non-polar matrices have failed94. Irradiation of crystals of ionic halides does lead formally to the production of trapped halogen atoms, but the properties of the active sites, the socalled "V-centres", leave little doubt that the trapped atoms interact strongly with a neighbouring halide ion to give a unit more aptly described as a [Hal-Hal'] - molecular ion94. Physical Properties of Halogen Atoms The physical properties of chlorine, bromine and iodine atoms are of two kinds: (i) those which can be precisely identified with the isolated atom, and (ii) those which can be defined or determined for the atom only when it is in the combined state. In practice, category (i) includes the nuclear properties of the different isotopes (Table 8), which are almost immune to variations of chemical environment, and those properties, determined mostly by spectroscopic methods, associated with the different energy states of the atom and derived ions (Table 10(a)); category (ii) includes properties such as atomic radius and electronegativity, whose significance demands that the atom be involved in some form of chemical bonding (Table 10(b)). Nuclear Properties: Isotopes32*45*95 ~97 Nine well, or reasonably well, authenticated isotopes of chlorine are known, ranging in mass number from 32 to 40. Of these only 35Q and 37C1 are stable, with natural abundances of 75-77 and 24-23%, respectively; the remainder are radioactive, decaying (i) by positron emission, accompanied by electron capture in one case (32Q, 33C1, 34Q, 36C1), or (ii) by jS-emission (36C1,38d, 39Q, 40Q). Isomers of 34Q and 38Q have been described. Isotopes of bromine more-or-less well-defined range in mass number from 74 to 90, but, apart from the very small contribution from unstable isotopes produced in nature by spontaneous fission and nuclear reactions induced by cosmic radiation, naturally occurring bromine consists of a mixture of the two stable isotopes 79Br (50-54%) and 81Br (49-46%). On current evidence, no less than four isobaric pairs of nuclei (mass numbers 77, 79, 80 and 82) have been characterized. Isotopes of bromine with mass numbers 74-78 and 80 decay by simultaneous positron-emission and electron capture, those with mass numbers 80 and 82-90 by βemission. Naturally occurring iodine is effectively mononuclidic containing the sole, stable isotope 127I. In all, 23 isotopes have been recorded with mass numbers between 117 and 139. In accordance with the usual pattern of behaviour, the lighter (neutron-deficient) radioactive isotopes with mass numbers 117-126 and 128, having a nuclear charge:mass ratio larger than is compatible with stability, undergo positron-emission or electron capture, both processes commonly occurring competitively and each resulting in the reduction of the nuclear charge by one unit. By contrast, the heavier (neutron-rich) isotopes with mass numbers 128 and 130-139, characterized by a nuclear charge:mass ratio which is too small for stability, attempt to redress the balance by ^-emission; the three heaviest, and 93 A. Carrington and D. H. Levy, / . Phys. Chem. 71 (1967) 2; A. Carrington, D. H. Levy and T. A. Miller, / . Chem. Phys. 47 (1967) 3801. 94 P. W. Atkins and M. C. R. Symons, The Structure of Inorganic Radicals, Elsevier (1967); K. D . J. Root and M. T. Rogers, Spectroscopy in Inorganic Chemistry (ed. C. N. R. Rao and J. R. Ferraro), Vol. II, p. 115. Academic Press, New York and London (1971). 95 C. M. Lederer, J. M. Hollander and I. Perlman, Table of Isotopes, 6th edn., Wiley, New York (1968). 96 G. H. Fuller and V. W. Cohen, Nuclear Data Tables, 5A (1969) 433. 97 Handbook of Chemistry and Physics, The Chemical Rubber Co. (1971-2).

1149

ATOMIC CHLORINE, BROMINE AND IODINE 137

138

139

shortest-lived, nuclei, I, I and I, also suffer neutron-emission. The atomic weights recommended by the IUPAC Atomic Weights Commission (1969)98 are set out in Table 10(a). Radioactive isotopes of the halogens are produced32»45»95 by the irradiation of stable isotopes of various elements with neutrons, with charged particles (protons, deuterons or α-particles) or with high-energy photons; bombardment by heavier projectiles, like 12C or 14 N, has yielded certain isotopes, access to which is otherwise difficult. The heavier isotopes of bromine and iodine are formed in the fission and spallation-fission of heavier elements. Indeed, isotopes of iodine and, to a lesser extent, of bromine played a prominent part in the early investigations of the fission of uranium and thorium, largely because these two elements were obtained in comparatively high yields and were relatively easy to isolate chemically and to purify; a number of radioactive isotopes of bromine and iodine were thus soon recognized as members of the decay chains which result from the fission process32. Neutron-irradiation now affords the commonest route to those isotopes of chlorine and bromine most widely used in radiochemical studies, viz. 36C1, 38C1, 80Br, 8omßr a n ( j 82ßr. In each case a stable nucleus absorbs a neutron and de-excitation of the highly excited composite nucleus thereby formed occurs by y-ray emission. Of the iodine isotopes of radiochemical significance, 128I is again the outcome of an η,γ reaction (involving naturally occurring 127I), while 131I is produced by the process 130Te -^L-> 131mje

^ 131 j

Z

e

>

1311

or by fission of uranium; more recently favoured in radiochemical studies, 125I and 132I are prodUCed thuS:

n>y

i24Xe —► i25Xe

Electron-capture tj = 17hr

β

► 1251

(ref. 99)

_

i32Te (uranium fission product)

> ™n

(ref. 100)

t | = 78 hr

It is a characteristic of nuclear reactions of the «,y-type that the target and product nuclei are chemically identical, a consideration which must limit the specific activity of the chemical product32. In other types of nuclear reaction, e.g. «,/?, «,α and fission, the nuclear reactant and product are chemically different, making possible, in principle, chemical separation to give so-called "carrier-free" active material. If the product of an «,y-reaction is itself short-lived with respect to a decay process in which there is a change of atomic number (e.g. β~, β+ or electron capture) giving a radioactive halogen nuclide, chemical separation of the product nuclide from the target is again feasible: such a process is exempli­ fied by the production of 125I from 124Xe and of 131I from 130Te. Although some nuclear reactions are thus better suited than others to the production of material of high specific activity, for a variety of reasons the material is seldom truly carrier-free. If a higher specific activity is required in the product of a simple «,y-reaction than that permitted by the available neutron flux, recourse may be made to the Szilard-Chalmers process. This depends on the fact that the recoil energy of the product nucleus of an η,γreaction (typically > 100 eV) is usually far in excess of ordinary chemical bond energies (1-5 eV), with the result that the nucleus recoils from its immediate chemical environment and may appear in quite a different chemical form. The phenomenon was first observed in 1934 by Szilard and Chalmers, who noted that, when ethyl iodide is irradiated with neutrons, 95.

98 Table of Atomic Weights, 1969. IUPAC Commission on Atomic Weights: Pure Appl. Chem. 21 (1970)

99 p. V. Harper, W. D. Siemens, K. A. Lathrop and H. Endlich, / . Nucl. Med. 4 (1963) 277. wo Bro. C. Cummiskey, S.M., W. H. Hamill and R. R. Williams, jun.,/. Inorg. Nuclear Chem, 21 (1961) 205.

1150

CHLORINE, BROMINE, IODINE AND ASTATINE: A. J. DOWNS AND C. J. ADAMS

most of the resulting radioactive iodine (128I) could be extracted from the ethyl iodide by water101. Three conditions must be fulfilled if a useful degree of enrichment of radioactive material is to be realized: formation of the radioactive atom must be attended by its dis­ engagement from the chemical environment (usually a molecule or molecular ion) character­ istic of the target; the atom must neither recombine with the fragment from which it has separated nor rapidly interchange with inactive atoms in the target material; and a chemical method must be available for the separation of the target compound from the radioactive material in its new chemical form. Enrichment of radioactive halogen isotopes produced by neutron-capture reactions (e.g. 38C1, 80Br, 82ßr and 128I) has been successfully achieved by the neutron-irradiation of an organic halogen compound such as CC14, CH2X2, CH3X, C2H5X, C4H9X, C2H4CI2 or C6H5X (X = Cl, Br or I) and aqueous extraction of the corre­ sponding halide ion. The reactions following rupture of the halide molecule, which are essentially those of the organic radicals and the highly energetic halogen atom formed by recoil, determine to a large extent the efficiency of separation; by contrast, the eventual fate of the active atom is apparently little influenced by the initial recoil energy. Szilard-Chalmers' separations of halogens have also been carried out by neutron-irradiation of solid or dissolved chlorates, perchlorates, bromates, iodates and periodates, from which the active halogen can be removed as silver halide after addition of halide ion carrier. Quite apart from its importance as a method of isotopic enrichment, the SzilardChalmers' process provides a much exploited opportunity to study the chemical behaviour of the highly energetic atoms produced by nuclear recoil. Such atoms are known colloquially as "hot" atoms, and their chemical reactions make up the realm of "hot atom chemistry" (seep. 1169). For fuller details pertinent to specific radioactive isotopes of the halogens, in respect of the principles and practice of formation, separation, enrichment, handling, detection and estimation, the reader is referred to more comprehensive or specialised texts32»33»45'102. Concerning the preparation of labelled chlorine and iodine compounds details have been compiled102; synthetic methods appropriate to such bromine compounds have also been outlined45. An elegant and effective method of labelling compounds is to inject them on to a gas-chromatography column charged with a radioactive halogen compound. By this method virtually carrier-free propyl bromide103 and arsenic and germanium chlorides104 have been obtained. Halogen-labelled molecular chlorides, bromides and iodides can be prepared by heating in vacuo an element such as silicon, boron or aluminium with labelled silver or copper(I) halides105»106. Of the several methods which have been described for the separation or enrichment of stable isotopes107, the following have proved most effective for the chlorine isotopes 35C1 and 37C133: thermal diffusion of gaseous HC1; ultracentrifuge treatment of molecular chlorides; electromagnetic separation of CuCl+ ions in a calutron (a type of cyclotron); migration of chloride ions in solution under the influence of an applied electric field; fractionation of chloride ions by ion-exchange; fractional distillation, e.g. of Cl2 or HC1; isotopic exchange, e.g. between gaseous Cl2 or HC1 and Cl ~ in aqueous solution. Virtually 101 L. Szilard and T . A . Chalmers, Nature, 134 (1934) 462^494. R. H . Herber (ed.), Inorganic Isotopic Syntheses, Benjamin, N e w Y o r k (1962). 103 F . Schmidt-Bleek, G . Stöcklin a n d W . Herr, Angew. Chem. 72 (1960) 778. 104 j . T a d m o r , / . Inorg. Nuclear Chem. 23 (1961) 158. 105 K . H . Lieser a n d H . Elias, / . Inorg. Nuclear Chem. 2 3 (1961) 139. 106 K . H . Lieser, H . W . Kohlschütter, D . Maulbecker a n d H . Elias, Z. anorg. Chem. 313 (1961) 193. 107 p. S. Baker, Survey of Progress in Chemistry, 4 (1968) 69.

102

ATOMIC CHLORINE, BROMINE AND IODINE

1151

complete separation of H35C1 and H37C1 has been achieved using the thermal gas-diffusion principle. Similar methods have been applied to the naturally occurring bromine isotopes 79ßr and ^Br, which are probably best separated by the thermal diffusion of HBrioe, though varying degrees of enrichment have also been brought about in the gas centrifuge109, by electromagnetic methods110, and by electrolytic transport in fused zinc or lead bromides111. In addition to data about the» decay and neutron-capture characteristics of different isotopes, Table 8 also alludes to the magnetic and electrical properties, where these are known. It is evident that, unlike the 19F nucleus, all of the chemically important nuclei of chlorine, bromine and iodine are quadrupolar with nuclear spins > £. The magnitudes of the nuclear spin, the magnetic moment and the electric quadrupole moment determine, most significantly, the principal characteristics of each nucleus with respect to nmr, esr and nqr experiments, although one or more of these parameters may modulate, via second-order interactions, other types of spectroscopic transition, as in the fine structure of microwave spectra or in the hyperfine structure of the atomic spectra. All five of the naturally occurring isotopes of chlorine, bromine and iodine, that is, 35C1, 37C1, 79Br, sißr and 127I, have featured in nmr, esr and nqr measurements. The effect of the quadrupole moment is to provide an efficient mechanism of magnetic relaxation with the result that, in normal chemical environments, the nuclei are invariably characterized by relatively diffuse magnetic resonances, which compare unfavourably with the narrow lines exhibited by 19F (/ = J). Nevertheless, some nmr measurements have been made to explore variations of chemical shift and linewidth, notably in relation to the possible effects of solvation and association of ions in solution (see p. 1239). On the other hand, the very properties which make the nuclei relatively unfavourable for conventional nmr studies make them highly eligible for nqr experiments, which afford a means of investigating the interaction of such quadrupolar nuclei with intramolecular electric fields; nqr measurements have been made for a large number of chlorine, bromine and iodine compounds112. With respect to the esr spectra of paramagnetic halogen-containing systems, the simplest examples of which are the halogen atoms (see Table 10), the nuclear properties of the naturally occurring isotopes determine the hyperfine structure of the spectra, there being 2/+1 equally spaced lines for an individual nucleus with nuclear spin /. Hence, information has been derived, not only about the properties of the halogen nuclei, but about the number of such nuclei in a chemical aggregate and about the density and orbital occupation of the unpaired electrons. The significance of some of the nmr, esr and nqr data is Considered in the contexts of the general characteristics of halide species (Section 3) and of individual halogen compounds to which they are relevant. A feature not evident from the nuclear properties of Table 8 concerns the potential of the different isotopes with respect to the Mössbauer effect, for which the source is a radioactive isotope of reasonable half-life. By radioactive disintegration the isotope populates an excited level which decays to the ground state by emitting low-energy yradiation. The only halogen isotopes which fulfil the necessary conditions are 127I and 129 I: 108 H.-U. Hostettler and Kl. Clusius, Proc. Intern. Symposium on Isotope Separation, p. 419. Amsterdam, North-Holland Publishing Co. (1957); Z. Naturforsch. 12a (1957) 974. 109 R . F . Humphreys, Phys. Rev. 56 (1939) 684. no C. W. Sheridan, H. R. Gwinn and L. O. Love, U.S. At. Energy Comm. ORNL-3301 (1962); W. Zuk, Ann. Univ. Mariae Curie-Skhdowska, Lubin-Polonia, Sect. A A, 12 (1960) 1. i n A. Lundon and A. Lodding, Z . Naturforsch. 15a (1960) 320; A. E. Cameron, W. Herr, W. Herzog and A. Lundon, ibid. 11a (1956) 203. 112 See for example E. A. C. Lücken, Nuclear Quadrupole Coupling Constants, Academic Press, London and New York (1969); M. Kubo and D . Nakamura, Adv. Inorg. Chem. Radiochem. 8 (1966) 257; H . Sillescu, Physical Methods in Advanced Inorganic Chemistry (ed. H. A. O. Hill and P. Day), p. 434. Interscience (1968).

24-23

37

^Ατ-α-α, p 40Ar-y-/>

WAi-n-p

40

37Cl-fl-y

35Cl-tf-y

39

38

Tracer element

Tracer element

36

38m

75-77

[ 5-482

5-57

13-2

0-712 114

55-5 m 1-4 m

ß~

37-3 m

ß~

ß~

0-74 s

IT

7-5

3-44

4-91

STAB LE ISOTO PE

3-lXl05y

STAB LE ISOTO PE

j3\EC

ß-

1-56 s

ß+

34*01 decay

34

35

320 m

|3 + , IT

3iP-a-/f

34m

2-5 s

)3+

32S-J-/I MS-p-n

33

0-31 s

)3+(«)

Half-life

Mode of decay

nS-p-n

Principal source

32

% natural Mass abundance number use as radiotracer

~ 7 - 5 , —3-2

1-91-3-45

1-11-4-81

0-714 OS")

4-50

2-5-4-5

4-51

9-5

Particle energies (MeV)

Cl, atomic number 17

1460-5800

246-1520

1600-3760

660

Ann. rad.

Ann. rad. 640-4100

2900

Ann. rad. 2210-4770

y-radiation energies (keV)

0-430 ± 0 1 0 0 (38C1) 0005(38md)

100 ± 3 0

44±2

Thermal neutroncapture crosssection (b)

3/2

2

3/2

(Ä/27T)

Nuclear spin, /

+0-68411

+ 1-285

+0-82183

Magnetic moment, μ (n.m.)

TABLE 8. NUCLEAR PROPERTIES OF THE BETTER DEFINED ISOTOPES OF CHLORINE, BROMINE AND IODINE*

III

-0062

-0017

-0079

Electric quadrupole moment, Q ( e x 10-24)

4-2 m

~5 s 4-4 h 17-6 m

IT

IT

j3 + ,EC

KSe-p-γ 79Br-/f-if'

79Br-/i-y

Tracer element

Tracer element

80m

80

79Br-/*-y somßr decay

3-573

1-365

4-6

2-72

~6-8

1-871 201

0085

STAB LE ISOTO PE

79m

50-54

j3 + ,EC

75 As-a-/i "Scr-d-n 7 8Se-p-rt 77 Se-/?-y

78

79

IT

76Se-/?-y

77m

75As-a-2/i

6-4 m

57 h

0 + ,EC

77

Tracer element

161 h

j8 + ,EC

As-a-3/i

75

76

1-7 h

|3 + ,EC

65CU-12C-2/I ™Se-d-n ™Se-p-y

75

36 m

)3\EC

«Cu-i2C-3/i

74 I

0-866 0-70-2-05

1-937,2-52

0-361

1-2-3-6

0-3-1-70

4-7

Br, atomic number 35

616-1257

37,85

Ann. rad. 6141

87-861005-2*

Ann. rad. 358-44436-7*

Ann. rad. 112-59621*

Ann. rad. 640

2-6±0-2 (80mBr) 8-5(80ßr)

1

5

3/2

3/2

1

±0-514

+ 1-317

+ 2106

±0-548

±018

+0-71

+0-31

±0-25

**Ki-n-p

Fission Th, U,Pu

Fission Th, U,Pu

Fission Th, U,Pu

1

1

86

87

88

89

90

Fission Th, U,Pu

Fission U

85

β-,η

β-.η

β~,η

β',η

β-

β-

β-

87Rb-/!-a fission Th, U,Pu

84

β-

β-

1 IT

β-

ΜΒτ-η-γ

Half-life

Decay energy (MeV)

1-6 s

45 s

16 s

55 s

54 s

30 m

31-8 m

2-41 h

35-5 h

|

1

61

7-1

2-8

2-6, 8 0

2-8

0-89-4-8

0-395, 0-925

0-97 4-8

0-257, 0-440

1-659 2-357

3 092

0046

Particle energies (MeV)

760

1440-5200*

305 0 (85mKr)

270-3930*

32-521

92-32056*

46 (IT), 698-41474-8

y-radiation energies (keV)

Br, atomic number 35 (cont.)

STAB ILE ISOTO PE

Mode of decay

82Se-*-y 83Se-j8-

Tracer element

49-46

Principal source

83

82

82m

81

% natural Mass abundancei use as number radiotracer

Table 8 (cont.)

1 3-0±0-2 (82mBr) 0-26 (82Br)

Thermal neutroncapture crosssection (b)

5

3/2

Nuclear spin, / (A/2*)

±1-626

+ 2-270

Magnetic moment, μ (n.m.)

±0-70

1 +0-26

Electric quadrupole moment, Q (ex 10-24)

100

Tracer element

127

128

127I_n_y

126Te-/?-/I

fi­

1-8-3-1

4-14

2150 1-251

0-149

3-17

25 08 m

1-27 2-14

113-212

ß+ 1129 ß~ 0-3851-25

0-79-2-13

1-2

2-36

-1-4

2-1,4-0

5-5

5-6

7

STAB LE ISOTO PE

jS ,EC

+

13 d

ß-

i8 + ,EC

i23Sb-a-W i25Te-d-n

126

60 d

EC

i23Sb-a-2« i24Te-i/-/i 125 Xe decay

Tracer element

125

4-2 d

ß+,EC

121Sb-a-« i23Sb-a-3/2

Tracer element

124

13-3 h

EC

i2iSb-a-2«

123

3-5 m

j3+,EC

122

i2iSb-a-3w 122Te_p-/i

i2iSb-a-4/i

121

19 m

2-1 h

I+protons i2iSb-a-5w ΐ2οχ β decay

120

ß+,EC

14 m

ß+,EC

Pd+i4N 1+protons

119

)3+,EC

lm

1-3 h

1 +protons

118

)3 + , E C

|3 + ,EC

La+protons

117

I, atomic number 53

442-9-969-5

Ann. rad. 388-71420*

35-48

Ann. rad. 602-7-2740*

159-781-4*

Ann. rad. 560-3450*

Ann. rad. 213-740*

Ann. rad. 560-1520

Ann. rad. 260, 780

Ann. rad. 550-1150

Ann. rad. 160,340

6·2±0·2

900 ± 9 0

1

5/2

2

5/2

2

5/2

+2-808

+ 30

-0-79

-0-89

Spall.-fission Pb, U; fission U, Pu; i33Te decay

Spall.-fission U; fission Th, U,Pu

134

i32Te decay following fission

Tracer element

132

133

i30Te-w-y; spall.-fission Th, U; fission Th, U, Pu

Tracer element

131

130

133CS-//-0C

U0To-d-2n ^Ί^-ρ-η 1291-n-y

130m

j

Fission U

Principal source

129

% natural abundance use as Mass number radiotracer

Table 8 (cont.)

20-9 h

ß~

52 m

2-3 h

ß~

β'

8 070 d

12-3 h

ß-

β'

0-62-1-7

0189

Particle energies (MeV)

4-2

1-80

3-56

1-10-2-46*

0-7-1-27

0-72-2-12*

0-970 0-257-0-806

2-99

1-7x107 y

ß~

8-82 m

0189

Half-life

IT

Decay energy (MeV)

Mode of decay

135-42467-1*

151-11592-5*

147-102395-0*

80-164722-92*

419-1150

39-58

y-radiation energies (keV)

I, atomic number 53 (cont.)

-0-7

18±3

(1301)

9±1

(130ml)

19±2

Thermal neutroncapture crosssection (b)

7/2

4

7/2

5

7/2

(Ä/27T)

Nuclear spin, /

+2-84

±3-08

+ 2-74

+ 2-617

Magnetic moment, μ (n.m.)

-0-26

±008

-0-40

-0-55

Electric quadrupole moment, Q (ex 10-24)

Fission U, Pu

Fission U, Pu

Fission U

Fission U

136

137

138

139

β~,η

)5-,n

ß~,n

ß-

ß~

2s

5-9 s

23 s

83 s

6-7 h

70

~2-8

135-42465-9* 200-3200*

0-5-1-4

2-7-7-0

7/2

The major sources of data for this table are: 1. Handbook of Chemistry and Physics, 52nd edn., B-253, The Chemical Rubber Co. (1971-72). 2. C. M. Lederer, J. M. Hollander and I. Perlman, Table of Isotopes, 6th edn., Wiley, New York (1968). 3. G. H. Fuller and V. W. Cohen, Nuclear Data Tables, 5A (1969) 433. Column 3 Source: Refers to the nuclear features (target element, projectile and outgoing particle, in order) whereby the radioactive isotopes are formed, p = proton; n = neutron; a = a-particle; d = deuteron; γ — y- or X-rays; spall .-fission = high-energyfission(followed by symbol of target element). 4 Mode of decay: IT = isomeric transition; EC = orbital electron capture. 5 Half-life: s = seconds; m = minutes; h = hours; d = days; y = years. 7,8 Particle andy energies: Ann. rad. refers to the 511 006 keV photon associated with the annihilation of positrons in matter. * = numerous well-defined energies {Rubber Handbook) within the limits specified. Thermal neutron-capture cross-section: b = barns (10~24 cm2). 9 10 Nuclear spin: units of Α/2π. 11 Magnetic moment: units of the nuclear magneton (n.m.), with diamagnetic correction. 12 Electric quadrupole moment: units are barns (10~24 cm2).

a

Spall .-fission U;fissionTh, U,Pu

135

1158

CHLORINE, BROMINE, IODINE AND ASTATINE: A. J. DOWNS AND C. J. ADAMS Mössbauer isotope

y-ray energy keV

Half-life of Mössbauer transition 10-9 s

1271

57-60

2-68

1291

Γ27-72 \27-78

16-8

Parent isotope

Half-life of parent isotope days

[ i27Te

105

( 127Xe i29Te

36 33

The form of the Mössbauer spectrum is determined by the characteristics of the nuclear transition and of the y-emission produced thereby, the lifetime of the excited state, the nuclear spins of the excited and ground state, and the field gradient at the nucleus. Nuclear isomer shifts and quadrupole splittings due to both the Mössbauer isotopes of iodine have been measured for a number of solid iodine compounds, e.g. I 2 , IC1, IBr, I2C16, alkali-metal iodides, iodates and periodates113, and these parameters have been correlated with details of the electron density, bonding and local symmetries of the iodine atoms. Radioactive halogens were much used in chemical studies even before pile-produced nuclides became generally available; accordingly there is a relatively long history of isotopic exchange and other tracer studies of the courses taken by chemical reactions and of the exploitation of radioactive halogen isotopes in chemical analysis and in various biological studies114. The increased availability of radioactive isotopes in more recent years has made possible a notable extension and development of these activities. Of the various uses of radioactive isotopes of chlorine, bromine and iodine, those listed in Table 9 are probably the most significant. Electronic and Thermodynamic Properties of the Isolated Atoms Table 10(a) presents for the isolated chlorine, bromine and iodine atoms details relevant to the atomic weight, spectroscopic properties, wave functions, promotion energies, ioniza­ tion potentials, electron affinity, electronic ^-factor and thermodynamic properties. In addition to the references given in the table, there exist substantial reviews of the optical and X-ray spectra of the atoms as reported up to 195632. The separations of the components of the inverted doublet which forms the ground state of each atom imply the following values (in cm- 1 ) for the one-electron spin-own-orbit coupling constants (£)" F, 269; Cl, 587; Br, 2456; I, 5069. Corresponding with this sequence, whereas the spin-orbit coupling in fluorine is adequately described by the Russell-Saunders laws, the large magnetic interaction in iodine is compatible, not with simple L,S- but withyy-coupling. The absence of low-lying excited states is evinced by the relatively high one-electron promotion energies calculated from the mean energies of the appropriate electronic states; these values are sufficiently high to call in question the authenticity of the concept of valence states involving the promotion of one or two /^-electrons and the subsequent formation of hybrid orbitals. For a given stage of ionization, the ionization potentials decrease relatively smoothly in the sequence Cl > Br > I after the relatively dramatic decrease in passing from fluorine 113 N. N. Greenwood, Chem. in Britain, 3 (1967) 56; M. Pasternak, Symposia of the Faraday Soc. 1 (1967) 119; D. W. Hafemeister, Advances in Chemistry Series, 68 (1967) 126; R. H. Herber, Progress in Inorganic Chemistry, 8 (1967) 1; J. Danon, Physical Methods in Advanced Inorganic Chemistry (ed. H. A. O. Hill and P. Day), p. 380, Interscience (1968). 114 j . Kleinberg and G. A. Cowan, U.S. At. Energy Comm. NAS-NS 3005 (1960).

The progress of exchange between two atoms of the same halogen in different chemical environments is followed by "labelling" one of the species with a radioactive halogen isotope, and subse­ quently separating the two compounds and measuring the variations of activity. Labelled halogen atoms may serve as indicators with respect to the progress, mechanism or the nature of inter­ mediates or products of a chemical reaction.

Studies of exchange*»*»0 in the following systems: X2/X" in aqueous solution; RX/X" in various solvents; RX/X2; HX/X 2 ; X 0 3 - / X 2 ; RX/AIX3; POCl 3 /Et 4 NCl in MeCN; PC15/C12 in CC14; Ph 2 I 2 /I in50%alcohol; PtX 4 2 -/X~ and MX 6 2-/X" (M = Re, Os, Pt or Ir) in solution. Possible auto-ionization in molecular halides like POCI3 and AsCl 3 .

1. (a) Fundamental studies of halogenexchange reactions for information (i) about the incidence of exchange, (ii) about the kinetics of the process, (iii) about its mechanism and (iv) about the equivalence of atoms in a molecule, e.g. PCI5 or PI12I2. (b) Investigations of reactions other than those involving isotopic exchange.

Such measurements are significant in the interpretation of solution theory and of the mode of action of ion-exchange resins.

Such methods are appropriate to the determination of traces (in the order 1 ppm) of the halogens in various systems; neutron-activation analysis is facilitated for chlorine, bromine and iodine (but not fluorine) by the relatively large neutroncapture cross-sections of the naturally occurring isotopes.

Measurements of the self-diffusion of X", Br0 3 ~ or I 3 ~ m in solu­ tion, of X" in crystalline AgX (X = Cl or Br)n and of X" or Br0 3 ~ in ion-exchange resins ;*>* determination of residence times in liquid extraction columns;0 solubility measurements, e.g. of/7-chloroiodobenzene in ethylene gas, p and investigations of precipitation processes, e.g. of Agl q and AgBr*. Detection and estimation of halogens can be achieved (i) by isotope dilution or related methods: e.g. I" estimated via 131 I in con­ junction with solvent-extractionr or ion-exchange ; s such methods have also been used in the analysis of fission products for bro­ mine and iodine ;a (ii) by neutron-activation analysis, used, for example, to estimate the halide content of aqueous media,*'* diverse organic systems,*»11 zinc sulphide phosphors,* SiC>2-Al203 catalysts,v and biological material,*^ and to evaluate the content and isotopic abundance of halogens in materials such as meteor­ ites.3'

2. Investigations of diffusion phenomena and of the distribution of components between two phases.

3. Radiochemical methods of analysis

Tracer studies of reactions such as that of Br atoms with aromatic compounds,* the iodination of metals,6 Friedel-Crafts and related reactions,' chlorination of hydrocarbons by BuOCl,* the I atom-catalysed isomerization of di-iodoethylene,h the polymeri­ zation of vinyl compounds by iodine-substituted free radicals,1 the action of bromine-containing inhibitors on the emulsion polymerization of styrene;1 studies of the stereospecificity of a bromination-debromination sequence starting from 1-bromocyclohexene,k and of the formation and reactions of organic free radicals (by trapping with labelled iodine).1

Comments

Examples

Application

TABLE 9. APPLICATIONS OF RADIOACTIVE HALOGEN ISOTOPES

Examples

Comments

Radioisotopes of chlorine, bromine and iodine, in suitable chemical These applications depend, for the most part, not on the chemical properties of combination, have found uses in the detection of leaks in process the radioactive material but on physical streams,8 the location of liquid junctions in oil pipelines,*»* in properties such as solubility or adsorp­ studiesof theflowof liquids and gases,e.g. atmospheric motions,** tion. in the detection offlawsin the sheathing of telephone cables, *»* and in various hydrological investigations, e.g. tracing water movement in soils,** evaluating the recharge-loss balance of ground water,00 and studies of sewage dispersion and water pollution.dd

»Supplement to Mellofs Comprehensive Treatise on Inorganic and Theoretical Chemistry·, Supplement II, Part I, pp. 1013-1063, 1080-1091, Longmans, London (1956). * Z. E. Jolles (ed.), Bromine and its Compounds, pp. 425-462, 786-798, Benn, London (1966). 0 M. F. A. Dove and D. B. Sowerby, Halogen Chemistry (ed. V. Gutmann), Vol. 1, p. 41, Academic Press, London and New York (1967). d S. May, M. Roux, Buu-Hoi and R. Daudel, Compt. rend. 228 (1949) 1865; G. Gavoret, J. chim. Phys. 50 (1953) 183; P. B. D. De la Mare and J. T. Harvey, / . Chem. Soc. (1956) 36; (1957) 131. β H. Sugier, Nukleonika, 12 (1967) 723 (Chem. Abs. 68 (1968) 108367t). f R. M. Roberts and G. J. Fonken, Friedel-Crafts and Related Reactions (ed. G. A. Olah), Vol. 1, p. 821, Interscience, New York (1963). * A. A. Zavitsas, / . Org. Chem. 29 (1964) 3086. h R. M. Noyes, R. G. Dickinson and V. Schomaker, / . Amer. Chem. Soc. 67 (1945) 1319. 1 K. Ziegler, W. Deparade and H. Kühlhorn, Annalen, 567 (1950) 151. 1 E. J. Meehan, I. M. Kolthoff, N. Tamberg and C. L. Segal, / . Polymer Set. 24 (1957) 215.

5. Technological and industrial uses.

4. Physiological and biochemical applica­Radiohalogens have been used to study the transport and distribu­ Apart from their practical clinical impor­ tance, radioisotopes of the heavier tions. tion of halide in mammalian tissues, e.g. the thyroid gland, halogens have afforded, via tracer central nervous system and bladder, and, to a limited extent, also experiments, intriguing results concern­ to trace halogen species in plant physiology.*»* In this context ing the transport of ions in the tissues of numerous compounds of biological interest have been labelled living mammals. It has also been shown with radiohalogens so that their metabolic fate may be explored. 82 that iodine entering the thyroid gland as Examples include the use as indicators of an Br-labelled anal­ iodide ion is oxidized forming first ogue of DDT,*** i3iMabelled insulin and antisera,* 82Br-labelled 7 82 131 monoiodo- and then di-iodotyrosine, proteins and steroids and Br- or I-labelled growth regulators which suffers oxidative coupling to form such as 5-bromouracil and 2-iodo-3-nitrobenzoic acid.*·* Clinical uses—131I is used for the diagnosis and therapy of thyroid thyroxine. Of the major organic con­ stituents of thyroid tissue, di-iodotyro­ disorders. A method has been developed for the direct irradiation, sine and thyroxine, the latter is believed employing 82Br, of malignant tissue in the bladder.*»* to be the circulating thyroid hormone which governs the metabolic rate of the whole body.*

Application

Table 9 (cont.)

48 (1952) 812.

J. Nölting, Z. Physik. Chem. {Frankfurt), 38 (1963) 154; M. Haissinsky, Nuclear Chemistry and its Applications (trans. D. G. Tuck), p. 553, AddisonWesley (1964). ° See, for example, M. Kubin, Proc. Symp. Radioisotope Tracers Ind. Geophys., p. 529, Prague (1966) {Chem. Abs. 68 (1968) 60954c). p A. H. Ewald, Trans. Faraday Soc. 49 (1953) 1401. q K. Müller and S. Karajannis, / . Radioanal. Chem. 2 (1969) 359. r H. G. Richter, Analyt. Chem. 38 (1966) 772. 3 M. Lesigang and F. Hecht, Mikrochim. Acta, (1962) 327. * I. F. Yazikov, N. N. Rodin, M. A. Dembrovsky and V. G. Lambrev, / . Radioanal. Chem. 3 (1969) 11. u R. Malvano and S. Kwiecinski, / . Radioanal. Chem. 3 (1969) 257. v P. Bussiere, A. Laurent and E. Junod, / . Radioanal. Chem. 2 (1969) 211. w P. Schramel, / . Radioanal. Chem. 3 (1969) 29; R. A. Nadkarni and W. D. Ehmann, ibid. p. 175. x A. Wyttenbach, H. R. von Gunten and W. Scherte, Geochimica et Cosmochimica Acta, 29 (1965) 467, 475. y J. Saroff, R. E. Keenan, A. A. Sandberg and W. R. Slaunwhite, jun., Steroids, 10 (1967) 15. z U.S. Pat. 3,370,173 {Chem. Abs. 68 (1968) 92298h). *» B. Keisch, R. C. Koch, A. S. Levine, J. Roesmer and W. S. Winnowski, U.S. At. Energy Comm. NSEC-120 {Nucl. Sei. Abstr. 21 (1967) 14328). ** A. Hamid and B. P. Warkentin, Soil Sei. 104 (1967) 279; E. Wagiel and J. Szymanski, Pr. Inst. Naft. (1968) 13 {Chem. Abs. 70 (1969) 30558a). cc K. Ubell, Proc. Symp. Isotop. Hydrol., p. 521, Vienna (1966). dd G. E. Eden and R. Briggs, Proc. Symp. Isotop. Hydrol, p. 191, Vienna (1966).

k C. L. Stevens and J. A. Valicenti, / . Amer. Chem. Soc. 87 (1965) 838. 1 See, for example, G. R. Martin and H. C. Sutton, Trans. Faraday Soc. m K. G. Darrall and G. Oldham, / . Chem. Soc. {A) (1968) 2584. n

1162

CHLORINE, BROMINE, IODINE AND ASTATINE: A. J. DOWNS AND C. J. ADAMS TABLE 10. PROPERTIES OF CHLORINE, BROMINE AND IODINE ATOMS

(a) Isolated atoms Property

Chlorine

Bromine

Iodine

Atomic number

17

35

53

Mass number of naturally occurring isotopes Atomic weight a

35(75-77%) 37(24-23%) 35-453

79 (50-54%) 81 (49-46%) 79-904

127(100%)

Optical spectra X-ray spectra Wavefunctions for electronic ground state

Refs. b, c Refs. f, g, h Refs. i, j

Refs. b, d Refs. f, g Refs. i, j

Refs. b, e Refs. f, g Refs. i, j

Electronic configuration and term of ground state 2 P$I2 -> 2P\n promotion energy, c m - 1 (kcal) One-electron promotion energies, c m - 1 (kcal)* ns2np5 -> ns2np\n+ \)sl ns2np5 -> ns2np4(n+ \)pl ns2np5 -> ns2np4ndl Ionization potentials, eV (kcal)b

h h h U h

[Ne]3s23pS P3/2

[Ar]35 Pm

126-9045

[ΚΓ]4ί/105ί25/75 P3I2

2

2

2

882-36 (2-523) b

3,685 (10-54)b

7,602-7 (21-76)b

72,500 (207)b 83,900 (240) b -91,000 (258)c

64,000 (183)b 75,100 (215)b -79,000 (225)d

55,400 (159)b 65,100 (186)b -70,000(200)·

12-967 (299-0) 23-80(549) 39-90 (920) 53-5 (1234) 67-80(1564)

11-84(273) 21-6(498) 35-9 (828) 47-3 (1091) 59-7(1377)

10-451(241-0) 1909 (440)

Electron affinity at 298°K, eV (kcal)k

3-68(84-8)

3-43 (790)

313 (721)

Esr properties: ^-factor (theoretical value assuming Russell-Saunders coupling = 1-3341064)

1-3339231

l-333921 m

1-333995°

Δ / Γ for JX2(g) -* X(g) at 298°K (kcal)°

28-989

23 036

18058

Thermodynamic properties of atoms at 298°K (ref. o) AHf° (kcal) AG/° (kcal) S°(caldeg-i)

28-989 25170 39-454

26-730 19-690 41-803

25-517 16-780 43-182

114 1-95 2-8 3-2

1-33 2-15 2-5 5-6

(b) Bound atoms Single-bond covalent radius (Ä) k Van der Waals' radius (Ä) p Electronegativity*1 Polarizability, Ä 3 (ref. r) Diamagnetic susceptibility, x 10 6 cgs units per g atom8

0-99 1-80 30 2-3 -201

-30-6

-44-6

* Calculated from the appropriately weighted mean energies of all the terms of both the ground and excited states. a Based on Table of Atomic Weights, 1969 IUPAC Commission on Atomic Weights; Pure AppL Chem. 21 (1970) 95.

ATOMIC CHLORINE, BROMINE AND IODINE

1163

b

C. E. Moore, Atomic Energy Levels, Vols. I-III, National Bureau of Standards Circular 467, Washington (1949-58); L. J. Radziemski, jun., and V. Kaufman, /. Opt. Soc. Amer. 59 (1969) 424; R. E. Huffman, J. C. Larrabee and Y. Tanaka, J. Chem. Phys. 47 (1967) 856. c Gmelins Handbuch der Anorganischen Chemie, 8 Auflage, "Chlor", Teil A, p. 123 (1968). d J. L. Tech, / . Res. Nat. Bur. Stand. 67A (1963) 505. • C. C. Kiess and C. H. Corliss, ibid. 63A (1959) 1. f J. A. Bearden, X-ray Wavelengths, U.S. Atomic Energy Commission, NYO-10586, Oak Ridge, Tennessee (1964); J. A. Bearden, Rev. Mod. Phys. 39 (1967) 78; J. A. Bearden and A. F. Burr, ibid. p. 125. * A. E. Sandström, Experimental Methods of X-ray Spectroscopy: Ordinary Wavelengths, Handbuch der Physik, 30 (1957) 78. h Ref. c, p. 133. 1 Functions based on Hartree-Fock-Slater approximation, F. Herman and S. Skillman, Atomic Structure Calculations, Prentice-Hall, Englewood Cliffs (1963). i Hartree-Fock wavefunctions, C. Froese, J. Chem. Phys. 45 (1966) 1417; J. B. Mann, Contract W-7405eng-36, Dep. CFSTI; Nucl. Sei. Abstr. 22 (1968) 17345. k A. G. Sharpe, Halogen Chemistry (ed. V. Gutmann), Vol. 1, p. 1, Academic Press, London and New York (1967); R. S. Berry, Chem. Rev. 69 (1969) 533. 1 V. Beltran-Lopez and H. G. Robinson, Phys. Rev. 123 (1961) 161. m J. S. M. Harvey, R. A. Kamper and K. R. Lea, Proc. Phys. Soc. (London), 76 (1960) 979. n K. D. Bowers, R. A. Kamper and C. D. Lustig, ibid, B70 (1957) 1176. 0 National Bureau of Standards Technical Note 270-3, January 1968; Codata Bulletin, International Council of Scientific Unions, Committee on Data for Science and Technology (1970). p L. Pauling, The Nature of the Chemical Bond, 3rd edn., p. 260, Cornell University Press, Ithaca (1960). q F. A. Cotton and G. Wilkinson, Advanced Inorganic Chemistry, 2nd edn., p. 103, Interscience (1966). r R. T. Sanderson, Inorganic Chemistry, p. 54, Reinhold (1967). 8 A. Earnshaw, Introduction to Magnetochemistry, p. 6, Academic Press, London (1968).

to chlorine; in this and other respects the variations in ionization potential are normal. Although unquestionably high, the first ionization potentials (I{) of chlorine, bromine and iodine are nevertheless lower than that of hydrogen (13.6 eV), and I\ for iodine is not much greater than for some metals—in particular zinc (9-39 eV) and mercury (10*43 eV). Even the sum of thefirsttwo potentials for iodine (29-54 eV) is only slightly greater than that for mercury (29-18 eV). Since the removal of one electron from the valence shell of a halogen atom leaves four electrons in the /?-orbitals, the X + cation is unlikely to be substantially smaller than the corresponding atom115. It is presumably this size factor, rather than the magnitude of the first ionization potentials, that limits the occurrence of mononuclear halogen cations under normal chemical conditions. The electron affinities of chlorine, bromine and iodine represent values recently recom­ mended115; they are based on an analysis of the results of three distinct methods, viz. (i) determination from computed lattice energies and other quantities in the Born-Haber cycle, (ii) the direct study of the equilibrium X(g)+e ^ X"(g) at a hot tungsten filament, and (iii) measurements of photochemical electron-detachment from halide ions in shockheated vapours of alkali-metal halides. The data of Table 10 confirm that the sequence of electron affinities is F < Cl > Br > I. Recently attributed1^ to a destabilization energy amounting to ca. 26 kcal g atom -1 , which accompanies the interaction of a fluorine atom with an external electron, the anomalous position offluorineis not really very significant in the comparative chemistry of the halogens. In chemical situations, factors such as lattice energies and solvation energies, which are sensitive functions of the size of the ions produced, invariably outweigh the small differences in electron affinity, and tend to dictate the differ­ ences in chemical properties among the halogens. Table 10(a) also includes values at 298°K for the standard heats of formation, free energies and entropies of the atomic halogens, together with the enthalpies of the reaction us A. G. Sharpe, Halogen Chemistry (ed. V. Gutmann), Vol. 1, p. 1, Academic Press (1967). 116 P. Politzer, /. Amer. Chem. Soc. 91 (1969) 6235.

1164

CHLORINE, BROMINE, IODINE AND ASTATINE: A. J. DOWNS AND C. J. ADAMS

JX2(g) -> X(g); tables of the usual thermodynamic functions are to be found elsewhere67. Variations in the entropies of the gaseous atoms are almost entirely determined by the term 3/2R In M+R In Q in the Sackur-Tetrode equation, M being the atomic weight and Q the electronic multiplicity, i.e. (2/+1), of the ground state. As with the molecular halogens and analogous molecular halides in the vapour phase, variations in the entropy terms are comparatively slight. The principal thermodynamic differences between the halogens arise from enthalpy effects which reflect the different strengths of the X-X bonds, viz. F-F < Cl-Cl > Br-Br > I-I. Several quite different interpretations of this sequence have been advanced, based, for example, on electronic repulsion, which is larger in F2 than in the other diatomic molecules (in keeping with the destabilization effect116 already referred to), or on partial multiple bonding in the heavier halogen molecules depending on the use, denied to fluorine, of valence-shell d-orbitals. One analysis117 concludes that electron repulsion is certainly a very important factor, but that the magnitudes of electron-nucleus attraction and nucleus-nucleus repulsion must also be taken into account; since all of these are a function of the internuclear distance, this quantity is possibly at the root of the some­ what anomalous sequence of bond energies. Properties of the Bound Atoms When a halogen atom engages in chemical bonding, its valence electrons are no longer identifiable with localized atomic orbitals. The gross perturbation of the valence electrons by ligand-fields means that most properties of the bound atom are, to some degree, con­ ditioned by its environment. However, the core electrons, being relatively much less responsive to such effects, may be justifiably regarded as occupying individual atomic orbitals. The photoelectron spectrum observed when a halogen atom in a compound is irradiated with X-rays furnishes energies of these core electrons, and it has thus been found, for example, that the energies of the ΑΓ-electron level and L\ sub-level vary almost linearly with the oxidation number of chlorine, the total shift being about 9-6 eV as the oxidation number varies from —1 to + 7118. The interatomic distances in the diatomic molecules of the gaseous halogens are conven­ tionally taken to be twice the single-bond covalent radii of the elements, which are set out in Table 10(b). With the uncertainty about the relative contributions of features such as electronic repulsion and ^/-orbital involvement in the bonding of these molecules, it is by no means clear that this step is sound, but no alternative is available at the present time. From the distances of closest approach of non-bonded halogen atoms, van der Waals' radii have been assigned by Pauling119; understandably the van der Waals' radius is much larger than the covalent radius of the halogen atom, being comparable with the ionic radius of the corre­ sponding halide ion. However, van der Waals' radii depend not only on the strength of the attractive forces holding the molecular aggregates together in the crystal, but also on the orientation relative to the covalent bond or bonds formed by the atom. Accordingly undue weight must not be given to the absolute magnitudes of the van der Waals' radii, which, even more than covalent and ionic radii, are the outcome of a highly simplified idea, The assessment of electronegativity or "the power of an atom in a molecule to attract electrons to itself"119 remains controversial principally because (i) no element has a unique electronegativity which remains constant throughout the whole range of its compounds, 117 G. L. Caldow and C. A. Coulson, Trans. Faraday Soc. 58 (1962) 633. us A. Fahlman, R. Carlsson and K. Siegbahn, Arkiv. Kemi, 25 (1966) 301. 119 L. Pauling, The Nature of the Chemical Bond, 3rd edn., Cornell University Press, Ithaca (1960).

ATOMIC CHLORINE, BROMINE AND IODINE

1165

and(ii) the effects of electronegativity cannot be completely extricated from those due to other bonding features. In view of these severe limitations, the electronegativity values given in Table 10(b) are based only on thermochemical data (the Mulliken and Pauling scales) or on estimates of effective nuclear charge and covalent radius (the Allred-Rochow formulation) 12<>; no reference has been made to the results of empirical or semi-empirical correlations of electronegativity with dipole moment, chemical shift, nuclear quadrupole coupling constant or vibrational properties of molecules. The electronegativity values are probably meaningful to no better than ±0-1 unit, serving only as rough guides, perhaps as the median numbers in a range for each element. By means of Pauling's expression £ ( A - X ) = V[B(A-A)

B(X-X)]+23(XA

- Xx)2

the energy of a bond between an element A and halogen X, B(A-X), may be estimated roughly in relation to the electronegativities XA and Χχ and to the energies of the homonuclear units A-A and X-X. Such estimates, the reliability of which has been tested for some forty halogen-containing bonds121, are useful as a guide, particularly when experimental data are lacking. The dipole polarizability, a, gives a measure of the susceptibility to deformation of the electronic charge cloud of an atom under the influence of an externally applied electric field. For a free halogen atom the quantity is amenable neither to experimental determination nor, with the present quality of knowledge about wave functions, to precise calculation122. For atoms in chemical combination, however, the molar refractivity of the system provides a means of estimating atomic polarizabilities; since halogen atoms are not commonly found in isotropic environments, values such as those given in Table 10(b) must be regarded as averages of the various components of the atomic polarizability tensor. The polarizability increases in unison with the total number of electrons in the series F < Cl < Br < I; it also increases with the single-bond covalent radius r in accordance with the general empirical rule123 a « 2-3r3. As might therefore be expected, the polarizability is smaller for the atom than for the corresponding anion122 and smaller for a halogen than for preceding atoms in a given row of the Periodic Table124. The principal importance of polarizability as an atomic property is in relation to intermolecular binding, notably through the medium of dispersion and Debye interactions, conspicuous, for example, in their influence on the melting and boiling points of molecular halogen compounds, which normally vary thus: AF% < ACU < ABr„ < AIW43. Chemical Properties of Halogen Atoms125 Although, directly or indirectly, a considerable body of information has been accumu­ lated, no systematic account of the chemical behaviour of the halogen atoms, as distinct from the molecules, has so far been given. In part this reflects the rather heterogeneous nature of experimental enquiries into reactions of the atomic halogens, which have been explored mainly with an eye to their mechanistic interest. Being produced under much less 120 F. A . Cotton and G. Wilkinson, Advanced Inorganic Chemistry, 2nd edn., Interscience (1966). 121 D . A . Johnson, Some Thermodynamic Aspects of Inorganic Chemistry, pp. 158-160 and 167-169, Cambridge (1968). 122 A . Dalgarno, Adv. Phys. 11 (1962) 310. 123 M. Atoji, / . Chem. Phys. 25 (1956) 174. 124 R. T. Sanderson, Inorganic Chemistry, p. 54. Reinhold, N e w York (1967). 125 G. C. Fettis and J. H . Knox, Progress in Reaction Kinetics, Vol. 2 (ed. G. Porter), p. 1, Pergamon, Oxford (1964); J. G. Calvert and J. N . Pitts, Photochemistry, p. 184, Wiley, N e w York (1966); D . M. Golden and S. W. Benson, Chem. Rev. 69 (1969) 125.

1166

CHLORINE, BROMINE, IODINE AND ASTATINE: A. J. DOWNS AND C. J. ADAMS

forcing conditions than are demanded by atomic hydrogen, nitrogen or oxygen, halogen atoms present fewer technical problems to the study of their reactions; on the other hand, the enhancement in reactivity with respect to the corresponding molecules is less pronounced, and the range of reactions open to the atom but not to the molecule correspondingly more restricted than in the case of hydrogen, nitrogen or oxygen. As already noted, the free atoms are believed to be essential intermediates in many reactions of the halogens and their com­ pounds. Examples of such reactions are given on pp. 1146-7. Inasmuch as the fundamental dissociation reaction X 2 ^ 2X appears to be the precursor of many reactions of the halogen molecules, the reactivity of the atoms is closely associated with the apparent reactivity of the molecules. Thus the reactions characteristic of the molecules—that is, oxidation-reduction, addition and substitution—all have their counterparts in the reactions of the atoms. The thermodynamic properties of the free halogen atoms inevitably make for more favourable free energy and enthalpy balances than in reactions involving the corresponding molecules, while the near-zero activation energies characteristic of many reactions of the atoms contrast strikingly with the substantial activation energies which are the general rule for reactions of the molecules (dissociation of which is commonly implicated). Thus, numerous reactions thermodynamically unattractive to the molecular halogens are feasible for the free atoms, though the products may be short-lived under normal conditions; many other reactions, thermodynamically feasible but kinetically slow for the molecules, proceed very rapidly with the free atoms. Early reports of so-called "active chlorine", produced by the action of an electrical discharge or by irradiation with ultraviolet light, refer to its abnormal chemical reactivity in its direct combination with ozone to produce chlorine monoxide and with sulphur and tellurium to form chlorides32. Abnormal reactivity is also illustrated by the report76 that iodine atoms, produced photochemically, will unite in chloroform solution with olefins at temperatures as low as — 55°C. Addition A primary reaction of the atomic halogens is the recombination 2X -> X2, as previously noted (see p. 1145), a relatively slow process in the gas phase, though highly susceptible to surface-catalysis126. Other examples of addition reactions of halogen atoms are: X

-f \

=6 \

• X·

+

Y-

Gas.phMCor/χ / solution /l \ X Crystal or ^ χ γ _ ρ ς γ = v

solution

'

s a m e o r dif]ferent

halogen) X·

+

YZ

Gas phase

w

χ γ ζ

( χ γ ζ = same or

halogen) I

+

D

Solution—^ j D

(2)*U28 different (3) 32 · 129

( D = = a r o m a t i c hydrocarbon, amine, alcohol or alkyl halide) (4)i3o.i3i

Cl·

+

02

Cl·

+

NO

^ s phase GaSphaSC

»■

-C1CO

(7)134

Gas

.

.HgCl

(8)i3S

Cl

+

CO

cl.

+

Hg

Gas phase _^

Phase

m

# a o o

^QC[

(5)l32 (6)133

ATOMIC CHLORINE, BROMINE AND IODINE

1167

Halogenation of many unsaturated organic compounds commonly proceeds by a free radical mechanism61. Photochemical chlorination and bromination involve, for example, an atomic chain process such as Cl · + C 2 H 4

-> C2H4CI ·

C2H4CI · + Cl 2 - * C2H4CI2+Cl ·

Similar chains are set up, in the presence of peroxide catalysts, in the reaction of sulphuryl chloride or of hydrogen bromide with unsaturated compounds, e.g. (a) R-CO-O · + H-Br

-> R-CO-OH+Br ·

(b) Br · + M e - C H = C H 2

-► Me-*CH-CH 2 Br

(c) M e - C H - C H 2 B r + H - B r

-> M e - C H 2 - C H 2 B r + B r ·

As a result of the reversible addition of halogen atoms, isomerization of olefinic compounds may also be induced; thus, dimethyl maleate is transformed to dimethyl fumarate in the presence of bromine atoms136. The action of molecular oxygen as a powerful inhibitor of many reactions of atomic chlorine—photochemical addition no less than chain reactions with hydrogen, hydrocarbons or carbon monoxide—depends on the scavenging action of reactions such as (5). Addition of a halogen atom X to a halide ion Y ~ affords the paramagnetic molecular anion XY _ , which can be identified by its esr, optical or Raman spectrum94»12^. In this way, the formation of the XY~ anion has been established as a result of y-irradiation, pulse radiolysis or flash photolysis of aqueous solutions or of low-temperature glasses containing halide ions; the so-called "V-centres" produced by the irradiation of crystalline ionic halides are similarly attributed to species of the type XY -. Under the appropriate conditions, y-irradiation of aqueous halide systems has also been found to yield anions identified by their esr spectra as XOH~ (X = Cl, Br or 1)137. The corresponding addition to a diatomic halogen or interhalogen molecule gives a triatomic radical of the type XYZ (X, Y, Z = the same or different halogen atoms). The formation of such species is supported indirectly by quantum-mechanical calculations, and 1261. M . Campbell and B. A . Thrush, Ann. Rep. Chem. Soc. 6 2 (1965) 37; S. W. Benson and W. B . D e M o r e , Ann. Rev. Phys. Chem. 16 (1965) 399; J. A . Kerr, Ann. Rep. Chem. Soc. 64A (1967) 7 5 , 1 2 1 ; M . A . A . Clyne, ibid. 65A (1968) 168; J. K. K. Ip and G. Burns, J. Chem. Phys. 51 (1969) 3414. 127 j . 1. G . Cadogan, Royal Inst. of Chem. Lecture Series, N o . 6 (1961); J. A . Franklin, G . Huybrechts and C. Cillien, Trans. Faraday Soc. 6 5 (1969) 2 0 9 4 . 128 M . C . R. Symons and W. T. D o y l e , Quart. Rev. Chem. Soc. 14(1960) 6 2 ; M . Anbar and J. K. Thomas, / . Phys. Chem. 68 (1964) 3829; H . C . Sutton, G. E . A d a m s , J. W . B o a g and B . D . Michael, Pulse Radiolysis (ed. M. Ebert, J. P. Keene, A . J. Swallow and J . H . Baxendale), p. 6 1 , Academic Press, L o n d o n (1965); B. Ceröek, M . Ebert, C . W. Gilbert and A . J. Swallow, ibid. p . 8 3 ; J. K. T h o m a s , Trans. Faraday Soc. 61 (1965) 7 0 2 ; R. C . Catton and M . C. R. Symons, / . Chem. Soc. {A) (1969) 4 4 6 and references cited therein; M . Hass and D . L. Griscom, / . Chem. Phys. 51 (1969) 5185; J. H . Baxendale and P. L . T. Bevan, / . Chem. Soc. (A) (1969) 2 2 4 0 . 129 L. Y . N e l s o n and G. C . Pimentel, / . Chem. Phys. 47 (1967) 3 6 7 1 ; Y . T. Lee, P. R. LeBreton, J. D . McDonald and D . R. Herschbach, ibid. 51 (1969) 455; D . H. Boal and G. A. Ozin, ibid. 55 (1971) 3598. 130 T. A . Gover and G. Porter, Proc. Roy. Soc. A262 (1961) 4 7 6 ; R. L. Strong, / . Phys. Chem. 66 (1962) 2423. 131 R. L. Strong and J. Perano, / . Amer. Chem. Soc. 89 (1967) 2535; A . M . Halpern and K. Weiss, / . Phys. Chem. 7 2 (1968) 3863. 132 E. D . Morris, jun., and H . S. Johnston, / . Amer. Chem. Soc. 9 0 (1968) 1918. 133 M . A . A . Clyne and D . H . Stedman, Trans. Faraday Soc. 6 4 (1968) 2698. 134 T. C. Clark, M . A . A . Clyne and D . H . Stedman, Trans. Faraday Soc. 6 2 (1966) 3354. 135 D . G. Hörne, R. Gosavi and O. P. Strausz, / . Chem. Phys. 4 8 (1968) 4758. 136 C. Walling, Free Radicals in Solution, p . 302. Wiley (1957). 137 R. C . Catton and M . C. R. Symons, / . Chem. Soc. (A) ( 1 9 6 9 ) 4 4 6 ; I. Marov and M . C . R. Symons, ibid. ( 1 9 7 1 ) 2 0 1 .

C.I.C. VOL II—OO

1168

CHLORINE, BROMINE, IODINE AND ASTATINE: A. J. DOWNS AND C. J. ADAMS

some authors have concluded on the basis of kinetic studies that Cl3 plays a significant part in the recombination of chlorine atoms32»33»138, though the most recent investigations do not favour this mechanism133. More direct evidence of the triatomic radicals comes12^ from molecular-beam studies of atom-recombination reactions and from matrix-isolation, whereby the radicals Cl3 and Br3, identified by their vibrational spectra, have been trapped in the condensate formed at low temperatures following the action of a microwave discharge on a gaseous mixture of the halogen with a noble gas. Δ// 298 ° for the reaction C1-+C12 -> *C13 has been estimated to be ca. 4 kcal139. Flash photolysis has been used130»131 to characterize charge-transfer complexes of the halogen atoms and to show, inter alia, that the stability constant of the Iatom-0-xylene complex in solution is larger than that of the I2-tf-xylene complex. Atom Transfer Atom-transfer reactions involving halogen atoms include and

X+H2

-+HX+H·

(1)32,140

X · + RH

-> H X + R ·

(R = organic group) (2)*5.i4i

which represent the propagation stages of the homolytic reactions between the elementary halogens and either hydrogen or organic compounds RH. The energetics of such reactions are compared in Fig. 10, which illustrates the increasingly endothermic character of

0CH 2 +HCI

FIG. 10. Energetics of hydrogen atom-transfer reactions of the halogen atoms. atom-transfer in the series Cl, Br, I, a variation which reflects the relative bond strengths of the hydrides HC1, HBr, HI. Thus, in contrast with the reaction chains with 103-1Q<* steps exhibited by vapour-phase chlorination reactions, bromination and iodination are character­ ized by relatively short kinetic chains or by simple radical reactions at all but the highest 138 E . Hutton and M . Wright, Trans. Faraday Soc. 61 (1965) 78. 9 V . I. Vedeneyev, L . V . Gurvich, V. N . Kondrat'yev, V. A . Medvedev and Y e . L . Frankevich, Bond Energies, Ionization Potentials and Electron Affinities, pp. 7 7 , 1 2 9 . Edward Arnold, London (1966). 14
ATOMIC CHLORINE, BROMINE AND IODINE

1169

temperatures. Moreover, compared with chlorine, bromine and iodine atoms are much more selective in their action, and in most cases the reverse action, viz. R · + HX -> RH + X ·

(R = H or organic group)

plays a significant role. Intensive studies suggest that atomic chlorine combines with carbon monoxide according to the following chain mechanism32»61: (i) Cl · + C O + M -> · C1CO + M* (ii) ClCO+Cl 2 ->OCCl 2 +Cl·

(M is probably Cl2)

The reaction is slow since (a) the initial step (i) requires a three-body collision, and (b) the intermediate radical tends to dissociate: (iii)

ClCO->CO+Cl

on the one hand and to be decomposed by the atom-transfer reaction (iv) Cl· + C1CO -^ Cl 2 +CO

on the other. Other examples of atom-transfer reactions are: I* + RI X· + YZ

-> I · + RI* -> X Y + Z ·

(R = organic group)!« (X = Cl, YZ = Br2, ClBr, C1I; X = Br, YZ = Cl2, BrCl;82 X = Ι(2Ρΐ/2), YZ = Cl 2 , Br2, IC1, IBr*43)

Cl · 4- NOC1 -> · NO + CI271 Br · + · C10 2 -> · BrO + · C1032 Cl · + ClOO · -> Cl 2 +0 2 i44

Cl · + NC13 -* · NC1 2 +Cl 2 "5 Cl · + C1N3 -> · N 3 + Cl2"5 Cl · + N 2 H 4 -* · N 2 H 3 + HCl"*

Apart from the more conventional studies of atoms in their electronic ground states, the behaviour of electronically excited halogen (2P\/2) atoms has also been examined via kinetic flash spectroscopy143. It has thus been found that, in the presence of CI2, Br2, IC1 and IBr, the decay of excited iodine atoms is dominated by reactions with these molecules, whereas, in the presence of CH3I, quenching of the I(2Pi/2) is favoured. It has also been found that hydrogen atom abstraction from paraffins is much more rapid for I(2Pi/2) than for I(2P3/2) atoms. The feasibility of such reactions, which are endothermic for the 2Py2 state of iodine, reflects the fact that the 2P\/2 ->2P3/2 transition is forbidden, so that the 2P\/2 halogen atoms enjoy relatively long radiative lifetimes: F 830 sec, Cl 83 sec, Br 1-1 sec, 10-13 sec. Provided that a population inversion of the 2Pm and 2^3/2 states is achieved in some primary photolytic step, the lifetimes of the 2P\/2 halogen atoms are such as to sustain laser action143. Thus, stimulated emission has been observed during flash photolysis of certain alkyl or perfluoroalkyl iodides, though not of molecular iodine: RI Ι( 2 Ρι/2)+Μ7603 cm-i)

U.V.

►R+ipPua) ► Ι(2Ρ 3/2 )+2/ινα603 cm"i)

Stimulated emission implicating Br(2p1/2) has also been observed following the flash photolysis of IBr. Investigations of such laser systems are potentially important inter alia for the light they may shed on primary photochemical processes and on energy-transfer mechanisms. 142 A. F. Trotman-Dickenson, Free Radicals, p. 91. Methuen, London (1959). !43 R. J. Donovan, F. G. M. Hathorn and D. Husain, Trans. Faraday Soc. 64 (1968) 1228; / . Chem. Phys. 49 (1968) 953; D . Husain and R. J. Donovan, Advances in Photochemistry, Vol. 8 (ed. J. N.Pitts, jun., G. S. Hammond and W. A. Noyes, jun.), p. 1, Wiley-Interscience (1971). 144 M . A . A . Clyne, Ann. Rep. Chem. Soc. 65A (1968) 183. 145 T. c . Clark and M . A . A . Clyne, Trans. Faraday Soc. 6 5 (1969) 2994. 146 S. N . Foner and R. L. Hudson, / . Chem. Phys. 4 9 (1968) 3724.

1170

CHLORINE, BROMINE, IODINE AND ASTATINE: A. J. DOWNS AND C. J. ADAMS

Hot-Atom Reactions*2·4™*-*1*™-™ Accelerated or electronically excited halogen atoms or ions, produced for example by an (η,γ) process, differ notably in their reactions from atoms with normal thermal energies, produced typically by photolysis. As distinct from thermal processes, so-called "hot" reactions are characterized by the following general features: (i) products are thereby realized which are not formed by thermal reactions; (ii) the reactions are temperatureindependent; (iii) inhibition may be brought about by chemically inert agents which dispel the kinetic, vibrational or electronic energy of the excited species before it undergoes reactive collisions; (iv) the reactions are not affected by low concentrations of reactive species which serve as scavengers for conventional free radicals. The first of these features is prominently displayed in the reactions of halogen atoms with hydrocarbons or organic halides: whereas halogen atoms with normal energies react primarily to abstract hydrogen atoms thereby forming hydrogen halides, "hot" halogen atoms, by virtue of their relatively high kinetic energy, electronic excitation or charge, are able to displace hydrogen atoms, organic radicals or other halogens: e.g. RCH 2 Br+Br*

• RCHBr· +HBr* or RCH 2 - + BrBr* RCH 2 Br*+Br· •RCHBrBr*+H· - RBr* + BrCH2 ■ or CH 2 BrBr*+R ·

In such a system, radicals may also result from the displacement of two atoms from a mole­ cule: overall it is thus possible to account, at least qualitatively, for the complicated pattern of products formed in the reactions of "hot" bromine atoms with organic bromides151. The principal technique for generating and studying "hot" atoms is nuclear recoil, certain transformations, e.g. 35C1-^I>36C1, 8omßr ^80ßr or I27i^i28i, being capable of producing species of very high energy, which, under properly controlled conditions, become available as "hot" atoms. Although nuclear reactions are normally rare phenomena and nuclear recoil accordingly yields a very small number of "hot" atoms, reactions resulting in the chemical combination of these atoms can readily be traced by means of their radio­ activity. A substance, which on nuclear transformation yields the desired "hot" atom, is mixed with the compound with which the "hot" atom is to react; such a mixture must be homogeneous on the scale of the recoil range of the "hot" atom. The sample is exposed to the appropriate nuclear radiation for a period at once long enough to produce the required number of activated atoms, but sufficiently short to avoid appreciable macroscopic de­ composition. The products incorporating the "hot" atoms are separated from one another and from the reactant, e.g. by Chromatographie methods. The species initially produced by nuclear transformations usually take the form of very high energy ions. These lose energy in successive interactions with the medium and should become neutralized in the process. Eventually the atoms drop into the range of chemical energies (< 100 eV) and in further collisions may react and combine by a "hot" reaction. However, it is not always clear whether the ionfirstformed becomes neutralized by chargeexchange processes before it loses enough energy to combine chemically. Whereas there is no evidence to implicate ions in the interaction with methane of 80Br produced by the (η9γ) 147 J. E. Willard, Ann. Rev. Phys. Chem. 6 (1955) 141. 14» j . E . Willard, Nucleonics, 19, No. 10 (1961) 61. 149 Chemical Effects of Nuclear Transformations, Proceedings of a Symposium, Prague, I.A.E.A. (1961). 1501. G. Campbell, Adv. Inorg. Chem. Radiochem. 5 (1963) 135. 151 M. Milman, Radiochim. Ada, 1 (1962) 15.

ATOMIC CHLORINE, BROMINE AND IODINE 79

128

1171

127

process from Br, I from the neutron-irradiation of I combines with methane partly as a "hot" atom and partly as a ground-state or excited ion152. Many of the current ideas on the mechanism of "hot" atom reactions in liquids and solids derive from the discovery of the Szilard-Chalmers effect in liquid ethyl iodide1**! and from subsequent investigations of halogen interactions in the condensed phase. The very exten­ sive literature relating to such interactions is mainly concerned with attempts to sort out the various types of process that affect the fate of the recoil species. Since it has usually not been possible to define clearly and to separate experimentally these events, comparatively little definite information has so far emerged concerning the detailed mechanism of actual "hot" reactions in the condensed phase. Thus, although considerable effort has been expended on the study of hot-atom reactions in solid halogen-containing compounds, e.g. salts of oxy-anions of the halogens79»153, most of the interest has focused on the fate of the halogen following activation by radiative neutron-capture. Subsequent annealing, whether induced by thermal or by radiative means, is an important factor in the outcome of such experiments, the effect being in general to restore a part of the initially separable recoil atoms to the form of the parent compound. As with the annealing of radiation damage in solids, this process works upon the defects produced in a matrix of otherwise normal crystal as a result of a sudden nuclear impulse. An informative way of examining the defects consists of following the change in the chemical state of the halogen atom responsible for the defects as a function of the time and temperature of annealing; such experiments supplement studies of radiation damage, which have usually depended on changes in the physical properties of the irradiated solids. Many features of the reactions of "hot" halogens are quite similar to those of "hot" hydrogen, which have been interpreted by the so-called "impact model"80,8i# According to this model, the principal mechanism whereby a "hot" atom combines in the gas phase embodies a single-step abstraction reaction. The recoil atom is assumed to enter into stable combination only over a certain energy range; above the upper limit, collisions are too energetic to permit formation of stable products, while the lower limit is determined by the minimum energy required to bring about reaction. By contrast, for all but grazing collisions, energy considerations are not of prime importance in determining the choice of reaction path from the several that are energetically accessible. Instead, the actual course of the reaction is determined (i) by the impact and steric parameters, (ii) by inertial restrictions in relaxation processes required for certain reactions, and (iii) by inductive effects, as yet poorly under­ stood. Apart from differences in steric properties, the reactions of "hot" halogens differ from those of "hot" hydrogen in that, for a given energy, the larger size of the halogen atom makes not only for a lower velocity but also for longer-lived and less localized collisions. There will therefore be more opportunity for the kinetic energy of the incident halogen to spread through the molecule with which it collides, and because the available energy is thus diffused over many bonds, it is less likely that any given bond in the impact area will be broken. Correspondingly there is a greater probability of rupturing two bonds to form a radical incorporating the "hot" atom151. The radioisotope produced by thermal neutron irradiation of a compound in dilute solution or in the gas phase is usually obtained in a chemical form other than that of the capturing molecule. However, when pure liquids or solids are irradiated, an appreciable 152 E . P. Rack and A . A . Gordus, / . Phys. Chem. 6 5 (1961) 9 4 4 ; / . Chem. Phys. 3 4 (1961) 1855. 153 See, for example, N . Saito, F . A m b e and H . S a n o , Radiochim. Acta, 7 (1967) 131; G. E . B o y d and Q. V. Larson, / . Amer. Chem. Soc. 9 0 (1968) 5092.

1172

CHLORINE, BROMINE, IODINE AND ASTATINE: A. J. DOWNS AND C. J. ADAMS

fraction of the total activity is found in the parent compound, presumably as the result of a secondary re-entry process. To explain this phenomenon, Libby made the general postulate that interactions of "hot" atoms at high energies with individual atoms in a molecule (billiard-ball mechanism), or at lower energies with the whole molecule (epithermal mech­ anism), could create a radical and leave the "hot" atom with insufficient energy to escape from the cage formed by its chemical environment154»155. After further collisional deactivation within the cage, the "hot" atom then combines with the radical previously formed. The following systems are representative of the reactions of "hot" halogen atoms (X*) which have been investigated in some detail: hydrocarbons +Χ* 45 » 149-151 ; RX-fX* (R = organic group) 45 ' 149 -* 51 ; olefins +X* 156 . 2.6. PHYSICAL PROPERTIES OF THE MOLECULAR HALOGENS

General Characteristics: Thermodynamic Properties The principal physical properties of the molecular halogens are summarized in Table 11. At ordinary temperatures the elements range from a greenish-yellow gas in chlorine, through a dense mobile liquid with a dark red colour in bromine, to a black, crystalline solid with a slight metallic lustre in iodine. With increasing atomic number, and hence polarizability, there is a strengthening of the van der Waals' forces between the diatomic molecules. The molecules persist throughout the solid, liquid and gaseous phases, though, as discussed below, intermolecular binding in the solid state becomes increasingly important in the series Cl2 < Br2 < hi perhaps the most dramatic evidence of this trend is supplied by the observation that, at high pressures, iodine crystals assume the electrical characteristics of a metal157. In their diatomic nature the heavier halogens differ notably from the elements of Groups V and VI, only the lightest members of which are diatomic under conventional conditions. A reflection of this difference is found in the melting and boiling points of chlorine, bromine and iodine, which are much lower than those of the previous members of their respective periods, that is, sulphur, selenium and tellurium. Table 11 summarizes, inter alia, recommended values of the boiling points, melting points, vapour pressures and heat capacities of the elementary halogens. Selected thermo­ dynamic functions, based on the most recent compilations67,i58-i59> a r e also presented for each element in a given phase and for the transitions from one phase to another. Thermo­ dynamic functions have been evaluated for the individual halogens over a wide range of temperatures by applying statistical mechanics to careful analyses of the rotational and vibrational terms of the molecular spectra, with corrections for rotational stretching, vibrational anharmonicity and rotational-vibrational interaction16^. Where comparisons have been made, the values calculated in this way are in very good agreement with those derived via the third-law method from measurements of heat capacity; there is thus no evidence that the solids retain residual entropy at limiting low temperatures. The results of the table confirm the similarity of the standard entropies of the halogens in the vapour phase. 154 w . F. Libby, / . Amer. Chem. Soc. 62 (1940) 1930; ibid. 69 (1947) 2523. 155 M. S. Fox and W. F. Libby, J. Chem. Phys. 20 (1952) 487. 156 See for example C. M. Wai and F. S. Rowland, / . Amer. Chem. Soc. 91 (1969) 1053. 157 A. S. Balchan and H. G. Drickamer, / . Chem. Phys. 34 (1961) 1948. 158 National Bureau of Standards Circular 500, U.S. Government Printing Office, Washington (1952). 159 National Bureau of Standards Technical Note 270-3, U.S. Government Printing Office, Washington (1968); Codata Bulletin, International Council of Scientific Unions, Committee on Data for Science and Technology (1970). 160 W. H. Evans, T. R. Munson and D . D. Wagman, / . Res. Nat. Bur. Stand. 55 (1955) 147.

(ii) Recommended thermodynamic value, D°29B1 Thermodynamic properties ΔΗ? at 298°K (kcal)1 AGt° at 298°K (kcal)' S° at 298°K (cal/deg mol)J

Gaseous molecules Electronic ground state, configuration ajhru4vg4 Excitation energies to lowest-lying excited states, r^cm""1) 3IIltt«-iE,+ 3Π0+««-ΐΣ,+ Ground state properties Internuclear distance, r e (A) Vibrational frequency, «».(cm"1) Annarmonic vibrational constant, o>e*e(cm-i) Force constant, £e(mdyne/A) Dissociation energy: (i) Spectroscopically determined, Do01

Molecular weight*

Property

—

cm'1 19,997-25 ±0-3

0 0 53-290

57-978

2-5141

35Cl22-675e 3-225° kcal eV 57-175 2-4793 35C12

1-9881· 35C12 559-72e

— 7-388 0-751 58-640

46072

1-9978

79Br8ißr 1-172»»· 2-475»·» cm'1 kcal eV 1 -9706 79fir2 15,894-5 45 -444 ±0-4 15,896-6 45-451 l-97098*Br2 ±0-5

2-2809»·» 79ßr8iBr 324-24'·»

13,814° 15,902-51*

0,2^1^4^,3 σΜ1 ««— σ^ττ« 4 π,4

ag2nu*irg3 au1 '*-ag27ru4frg4

~ 23,500 (continuum)1» 17,672-6°

1Σ.+

159-808

Br 2

1Σ.+

70-906

Cl 2

TABLE 11. PHYSICAL PROPERTIES OF MOLECULAR CHLORINE, BROMINE AND IODINE

—

14-919 4-627 62-277

36115

1-5661

i27I 2 0-60738 h l-720 b 1 cm' kcal eV 12,440-9 35-570 1 -5424 « 7 i 2 ±1-2

2-666h i27i 2 214-52 h

11,888° 15,769-48d

ap-TTutiTflGu1 **— af-irutiTg*

1Σ,+

253-8090

I2

Intramolecular distance,
Crystal Structure Molecules per unit cell Unit cell dimensions (Ä)

Liquid

Critical pressure (atm) Critical density (g/cm3) Vapour pressure, p (mm Hg) Solid

,

2-715 ± 0 0 0 6

2-27±0-10 3-31-4-14

1-980 ± 0 0 1 4 3-32-3-97

3496-4-412

Orthorhombic, space group Cmca* 4 a b e 7-136 4-686 9-784 (1<0°K)

Orthorhombic, space group Cmcax 4 a b e 6-67 4-48 8-72 (123°K)

log/7 = - ( 3 5 9 4 0 3 / r ) + 0 0 0 0 4 4 3 4 r -2-9759 log T + 18-80572r'** log/7 = - ( 2 9 7 0 / D - 3 - 5 2 log T + 18-751*'**

~116 q l-64 p

°C 546p

Orthorhombic, space group Cmca1 4 a b e 6-24 4-48 8·26(113°Κ)

102n l-26 p

315p

log/7 = -(11310·00/:Τ)+0·17483:Γ - 1 8 4 1 7 5 log T +444-26938 r log/7 = -(2047·75/Γ)-0·006100Γ +0-9589 log Γ+8-65047'

76-1° 0-573°

588 p

27-758(c)J

36-379(l)J

53-290( g y 144-0°

14.477m'**

10-220m

6-965n

°K 819p

21-87m'ss 3-709m·**

21-24 m 2-527m

+ 113-60*

386-75m m

°C +185-24m'ss

eV 9-22 2-55

°K 458-39m'ss

20-40» l-531 n

[

kcal 212-6 58-8

100255 **

+59-47 m + 58-78* -7-25m

eV 10-51 2-55

7 064

332-62m 331 -93* 265-90 m

kcal 242-4 58-8

12

4-878"

-101001

-34051

eV 11-51 2-45

Br 2

m

Cl 2

log/? = -(2890·8/Γ)+0·09914Γ -58-836 log T + 132-26593' logp = -(1414·8/Γ)-0·01206Γ + 1-34Χ10-5Γ2 + 10-91635*

417-16°

172161

239101

Properties of the condensed phases Boiling point

Melting point Heat of vaporization at boiling point, A£Tvap(kcal/mol) Entropy of vaporization at boiling point (Trouton constant) (cal/deg mol) Heat of fusion (kcal/mol) Heat of sublimation at melting point (kcal/mol) Entropy of reference state at 298°K, 5°(cal/deg mol) Critical temperature

kcal 265-4 56-5

Adiabatic ionization potential, Χ2 + ΡΠ 3 / 2 .«>«-Χ2( 1 Σ. + ) , ί Adiabatic electron affinityk

Property

Table 11 (cont.)

Mechanical properties Density, i/(g/cm3) Solid

Liquid Gas Thermal conductivity (cal/sec/ cm2/°C/cm) Solid Liquid Gas Coeflacient of cubical expansion (xl04per°C) Solid Liquid Gas

Thermal properties Heat capacity, Cp(cal/mol deg) Solid

Twice van der Waals' radius ofX(A) „ .
w

8-616-9-562 (298-6000°K) m

3-04-2-52 X10-4 (283-323°K) z 1 -03-2-72 X10-5 (273-700°K) z

- 4 - 9 5 x 10-4 (298°K) y 1 -32-4-59 X10-5 (200-600°K) z

d=

2-098-3-5 Χ10-4Γ (77-4-158-2°K)i

4 1 7 - 4 0 5 (0-123°K) b d= 3-924-1-062x10-3/ (t= - 2 3 - 5 to -106°C) b b

2-47 (167-250°K) bb ll-0(273-323 o K) b > cc

18-579-18-077 (265-9-300°K)

16-03-15-70 (172-239°Κ)*Ί

7-576-9-710 (200-6000°K) m

1·68(77·4-158·2°Κ) ω 13-07-54-48 (182-383°K)i 36-61-38-33 (273-373°Κ,/> = 0-1 atm) 1

1-725-14-732 (15-265-9°K) w Cp=

5-15-4-886 (78-333°K)b»<*

2-81 (273-387°K) ce 8-54b

1 -20-1 -52 X10-5 (500-600°K) z

1-065x10-3 (297-316°K) b

-121098+0059012Γ + 6-686χ10 5 Γ" 2 (10-330°K)XSS 19-281 (387-458°K) m 8-814-9-746 (298-6000°K) m

2156(253°K)V 0-175*

765-86 (4°K)V 0-20 v

108-95 (20°K)V ^008v

0-810-13-27 (14-05-172°K)l>.i

1-29

1-46

1-68

4-30

3-90

3-60

OS

-j

► —

Liquid Gas at N.T.P. Specific conductivity, k (ohm - 1 c m - 1 ) Solid

Electrical, optical and magnetic properties Dielectric constant, e Solid

Gas Surface tension of liquid, y(dyne/cm)

Viscosity, ^(centipoise) Liquid

Compressibility ( x 10^ per bar) Solid Liquid

Density, d (g/cm ) (cont.) Liquid Gas

3

Property

Table 11 (cont.)

d=

—ψ

[1 +

T2

)

P +

2-8-8-1 X 10-13 (253-266°K)**

3-255-3-080 (273-303°K) dd 10047 b

γ = 45-5-0182t (over liquid range, t in °C)Be

31-2-18-4 (213-293°K)i

2-142-1-947 (203-283°K)i 100152 b

η= 1-241/(1+0-012257/ +2-721x10-6/2) ( / = - 4 - 3 to + 32°C)b 001526-004292 (293-873°K) b

—

62-5-49-0 (0-500 bar, 293°K)ff

3-187-3-009 (273-325°K)b>cc»dd 0005480-0005038 (361-386°Κ,/> = 1 atm) b

Br 2

77 = 0-385/(1+0-005878/ -0-00000392/2) (/ = - 7 6 - 5 to -33-8 0 C) to 001294-003209 (289-772°K)i

—

116-83 (0-500 bar, 293°K) ff 88-9-47-4 (180-240°K)°

/2515-2·991(Γ+199)\2 2 l Γ2 ) ^ 2 + ···1 (288-348°K, p up to 2 atm) b

I

/2515-2·991(Γ+199)\

1-6552-0-5672 (203^17°K) b . 0-86427/. n^

Cl 2

1-71-12-0x10-7 (298-383°K) cc Single crystal at room temperature: 5x10-12 normal \ 1 7 x 1 0 - 8 parallel )to^PlaneV 0-8-l-7xl04(350kbar) 3 J

10-3 (296°K)CC Single crystal at ~300°K: ee = 6;e b = 3 ; e c = 40 i i 11-08-12-98 (391-441°K) b
36-88-34-04 (398-428°K) hb

001785-003604 (379-796°K) b

2-268-1 -414 (389-458°K) cc

—

10-1-25 (0-200 kbar)ee

3-960-3-736 (393-453°K)CC 0011320-0007434 at N.T.P. (723-1274°K, p = 0-04-1-05 atm)q

Ϊ2

+ 1-356"

-40-2 ~-35qq + 1065"

-53-6 ~-70b

+0-535"

Single crystal: Xa = - 8 4 - 4 ; Xb= -89-3 ;X e = -75-9°° - 9 9 to -84(391-432°K) k k —

pp

~-60b

— 1

28-32 m m

17-61 mm

11.74mm

50-9 x 10-6 (394-413°K) b ' cc nD = 3-34(293°K) kk nD= l-98 k k λ = 6708-5000 Ä n = 1 00197-1 00212, exhibits anomalous dispersion1111

log& = -ll-372+0-0128t (over liquid range, / in °C)eg — nD= 1-6083 (293°K) mm λ = 6708-5461 Ä n = 1 0011525-1-0011849nn

— nD= 1-3788 (298°K) n λ = 6708-4800 Ä n = 1 00077563-1 00079166nn

<1χ10-ΐ6ΐ

c

b

Table of Atomic Weights, 1969, IUPAC Commission on Atomic Weights; Pure AppL Chem. 21 (1970) 95. Supplement to Mellows Comprehensive Treatise on Inorganic and Theoretical Chemistry, Supplement II, Part I, Longmans, London (1956). G. Herzberg, Molecular Spectra and Molecular Structure. I. Spectra of Diatomic Molecules, 2nd edn., pp. 512-541. Van Nostrand (1950); W. G. Richards and R. F. Barrow, Proc. Chem. Soc. (1962) 297. d J. I. Steinfeld, J. D. Campbell and N. A. Weiss, / . Mol. Spectroscopy, 29 (1969) 204; J. A. Coxon, ibid. 37 (1971) 39. 6 A. E. Douglas, Chr. Kn. Möller and B. P. StoichefF, Canad. J. Phys. 41 (1963) 1174; M. A. A. Clyne and J. A. Coxon, / . Mol. Spectroscopy, 33 (1970) 381. f J. A. Horsley and R. F. Barrow, Trans. Faraday Soc. 63 (1967) 32. * R. J. LeRoy and G. Burns, / . Mol. Spectroscopy, 25 (1968) 77. h D. H. Rank and B. S. Rao, / . Mol. Spectroscopy, 13 (1964) 34. 1 R. J. LeRoy and R. B. Bernstein, / . Mol. Spectroscopy, 37 (1971) 109. J National Bureau of Standards Technical Note 270-3, January 1968; Codata Bulletin, International Council of Scientific Unions, Committee on Data for Science and Technology (1970). k D . C. Frost, C. A. McDowell and D. A. Vroom, J. Chem. Phys. 46 (1967) 4255; A. W. Potts and W. C. Price, Trans. Faraday Soc. 67 (1971) 1242; A. P. M. Baede, Physica, 59 (1972) 541. 1 Gmelins Handbuch der Anorganischen Chemie, 8 Auflage, "Chlor", Teil A (1968).

a

Standard potential for the system iX2+e-X-(aq),£°(V)

Liquid Gas

Molar refraction (cm3/mol) D-line, various temperatures Diamagnetic susceptibility, χ(106 cgs units/mol) Solid

Refractive index, n Solid Liquid Gas at N.T.P.

Liquid

Landolt-Börnstein Tables, II Band, 2 Teil, Bandteil a, Springer-Verlag (I960). * J. Donohue and S. H. Goodman, Acta Cryst. 18 (1965) 568. u F. van Bolhuis, P. B. Koster and T. Migchelsen, Acta Cryst. 23 (1967) 90. v R. Bersohn, /. Chem. Phys. 36 (1962) 3445; E. A. C. Lücken, Nuclear Quadrupole Coupling Constants, Academic Press, London and New York (1969); N. Nakamura and H. Chihara, J. Phys.Soc. Japan, 22 (1967) 201. w D. L. Hildenbrand, W. R. Kramer, R. A. McDonald and D. R. Stull, /. Amer. Chem. Soc. 80 (1958) 4129. x D. A. Shirley and W. F. Giauque, / . Amer. Chem. Soc. 81 (1959) 4778. y The Encyclopedia of the Chemical Elements (ed. C. A. Hampel), Reinhold, New York (1968). z Landolt-Börnstein Tables, II Band, 5 Teil, Bandteil b, Springer-Verlag (1968). *» L. L. Hawes and G. H. Cheesman, Acta Cryst. 12 (1959) 477. bb L. L. Hawes, Acta Cryst. 12 (1959) 34. cc Kirk-Othmer's Encyclopedia of Chemical Technology, 2nd edn., Interscience (1963-9). dd G. Fröhlich and W. Jost, Chem. Ber. 86 (1953) 1184. θβ R. Grover, A. S. Kusubov and H. D. Stromberg, /. Chem. Phys. 47 (1967) 4398. » T. W. Richards and W. N. Stull, /. Amer. Chem. Soc. 26 (1904) 408. 99 M. S. Chao and V. A. Stenger, Talanta, 11 (1964) 271. hh S. Kim and S. Chang, Daehan Hwahak Hwoejee, 9 (1965) 110 (Chem. Abs. 64 (1966) 5785c). 11 M. Simhony, /. Phys. Chem. Solids, 24 (1963) 1297. Ji A. S. Balchan and H. G. Drickamer, / . Chem. Phys. 34 (1961) 1948. kk Gmelins Handbuch der Anorganischen Chemie, 8 Auflage, "Iod" (1933). 11 A. W. Francis, /. Chem. Engineering Data, 5 (I960) 534. mm P. M. Christopher, Austral. J. Chem. 20 (1967) 2299. nn C. Cuthbertson and M. Cuthbertson, Phil. Trans. 213 (1913) 1. 00 J. De Villepin, /. chim. Phys. 59 (1962) 901. pp S. Broersma, / . Chem. Phys. 17 (1949) 873. qq P. Pascal, Nouveau Traiti de Chimie Minerale, Vol. XVI, Masson et Cie, Paris (1960). ™ A. G. Sharpe, Halogen Chemistry (ed. V. Gutmann), Vol. 1, p. 12, Academic Press, London and New York (1967). ss For more recent values see B. Lindenberg, Compt. rend. 273C (1971) 1017.

m JANAF Thermochemical Tables, The Dow Chemical Company, Midland, Michigan (1960-8). n D. R. Stull and G. C. Sinke, Advances in Chemistry Series, 18 (1956). 0 T. R. Thomson, H. Eyring and T. Ree, Proc. Nat. Acad. Sei. 46 (1960) 336. p J. A. Burriel Lluna, C. B. Cragg and J. S. Rowlinson, An. Quim. 64 (1968) 1. q M. L. Perlman and G. K. Rollefson, / . Chem. Phys. 9 (1941) 362. r A. N. Nesmeyanov, Vapor Pressure of the Chemical Elements (ed. R. Gary), Elsevier (1963). 8

PHYSICAL PROPERTIES OF THE MOLECULAR HALOGENS

1179

The Gaseous Molecules The bonding in the diatomic halogen molecules can be described approximately in terms of simple molecular-orbital theory (see Fig. 11) as depending on the balance between six electrons in bonding orbitals ( a ^ ^ a n d four in anti-bonding orbitals (π^4). This represents a

\

\

\

>-H4K

\

/

/

N

?s ^

i \

/

\

I

y

J

\

\

/ ^

/

^

f

^

t

v

/

/

/

s \

/ . - - ■ '

"-H—\

;H-l·-"

FIG. 11. Schematic energy level diagram showing molecular orbitals for the diatomic halogen molecules.

diamagnetic ground state with a bond order of unity, though a more sophisticated descrip­ tion which takes account of interactions of the «rf-orbitals would impute at least some (p-d)7r-character to the X-X bond. The outer ng and nu orbitals are mainly atomic, being localized on the halogen atoms. The electronic spectra of the halogens have attracted much study161. The colours of the gaseous systems arise from absorption bands corresponding to the electronic transitions < > , 2 * « 4 W (3Π„ and HIU) +- σ , ζ , τ ^ ( i s , + )

in which an electron is excited from the anti-bonding wg to the anti-bonding au orbital. With increase of atomic number there is a decrease in the energy separation of these two orbitals, together with an enhanced probability for the singlet-triplet transition to the lower energy 3IIW state. These features together account for the variations of frequency and intensity exhibited by the visible absorption bands in the series of gaseous molecules CI2, Br2, I232,74,161,162. However, there is a complication in that the electronic states of Cl2, Br2 and 161 J. A. Coxon, Molecular Spectroscopy, Vol. 1 (ed. R. F. Barrow, D. A. Long and D. J. Millen), Chemical Society Specialist Periodical Report (1973), p. 177. 162 G. Herzberg, Molecular Spectra and Molecular Structure. /. Spectra of Diatomic Molecules, 2nd edn., vanNostrand (1950); R. W. B. Pearse and A. G. Gaydon, The Identification of Molecular Spectra, 3rd edn., Chapman and Hall, London (1963).

1180

CHLORINE, BROMINE, IODINE AND ASTATINE: A. J. DOWNS AND C. J. ADAMS

I2, unlike those of F2, approach Hund's coupling case c, wherein the spins and angular momenta of the individual electrons first combine to give separate resultants j \ 9 j 2 , etc., which then yield a final resultant denoted by the quantum number Ω. This situation, corresponding to (j,j) coupling in atoms, renders it impossible to assign a value of the conventional quantum number Λ or of the multiplicity; instead, detailed correlations of Ω with atomic J are necessary. Chlorine gas shows a strong continuous absorption extending from the blue to about 2500 Ä and having a maximum near 3300 Ä. With greater path-lengths, there is also seen a weak banded absorption spectrum sharply degraded to the red extending from about 5800 A to a convergence limit near 4785 Ä. Though modified in details of position, intensity and complexity, these characteristics of a continuum in the blue-ultraviolet region and a banded system close to 5000 Ä are also found in the absorption spectra of gaseous bromine and iodine. The transition 3n o+w (0 + w)<- ^g+ is responsible for the system of bands in the green, which shows a well-marked convergence to a continuum at ca. 20,880, 19,580 and 20,040 c m - 1 for Cl2, Br2 and I2 respectively, thresholds corresponding to the energy of the dissociation process Χ 2 ->Χ(2Ρ 3 /2) + Χ( 2 Λ/2)

Subtraction of the atomic energy of excitation of X(2P1/2) from the energy corresponding to the limit gives the most reliable values at present available for the normal dissociation energy of each of the X 2 molecules; the most recent extrapolation of the convergence limits for Cl2, Br2 and I2 forms the basis of the results in Table ll 1 6 3 . The circumstance of absorption in which the observable transitions lead, as the frequency increases, first to bands of higher vibrational quantum numbers and finally to continuous absorption clearly implies the disposition of potential energy curves illustrated in Fig. 12, the shallow minimum

Energy X(2pi/2> + X6P 3/2 ) X(2p3/2) +X(2p 3/2 )

Intemuclear distance FIG. 12. Schematic form of the potential energy curves of the lowest observed electronic states of the molecules Cb, Br2 and I2. 163 R . j . LeRoy and R. B. Bernstein, / . Mol. Spectroscopy, 37 (1971) 109.

PHYSICAL PROPERTIES OF THE MOLECULAR HALOGENS

1181

3

for the upper n 0 + w state being at an internuclear distance about 0-4 A greater than that of the ground state. Detailed vibrational analyses of the banded spectra of the gaseous halogens yield the values given in Tables 11 and 12 for the vibrational frequencies ωβ, the anharmonicity constants a>exe and (*>eye, and the stretching force constants ke. Where comparisons can be made, good agreement is found between these values and the results derived, for example, from the Raman spectra, as in recent studies of the resonance Raman effect and resonance fluorescence of the gaseous halogens164. The vibrational frequencies of liquid chlorine and bromine are not significantly different from those of the gaseous molecules, but the Raman spectra of crystalline chlorine and bromine at low temperatures165 disclose an appreciable reduction of the vibrational frequencies corresponding to a reduction in ke of 9% for Cl2 and 19% for Br2. A decrease in the X-X stretching frequency is also found commonly to attend the transition from the vapour to the solution phase. The significance of this and related spectroscopic observations is indicated subsequently in the general context of the behaviour of the halogens in solution (pp. 1198-1200). Detailed rotational analysis of individual vibrational bands of the electronic transition 3 n 0 + w <- χΣ8+ leads to the most reliable estimates of the rotational constant Be and hence of the internuclear distance re, associated not only with the ground state (ιΣ8+) but also with the 3 n 0 + w excited state of each molecule. Values of these constants and of the rotational constants ae and De, derived from the most recent of such analyses, are collected in Tables 11 and 12. Various attempts have been made to resolve the continua in the electronic spectra of the halogens into components attributable to distinct transitions161. Some of the excited states implicated in these transitions, seldom well-defined and commonly unstable with respect to dissociation, are listed in Table 12; in most cases positive experimental information is lacking and some of the details must therefore be regarded as no more than tentative. Through the action of heat or of an electric discharge on the gaseous halogen or through certain reactions which furnish excited atoms or molecules (e.g. that between hydrogen and the elementary halogen), emission spectra have also been observed; again, despite numerous studies, many features of these spectra have yet to be interpreted satisfactorily. All three halogens exhibit fluorescence to some degree32»161. In this respect iodine has been one of the most widely studied systems, fluorescence occurring in the green as well as in the ultraviolet between ca. 1800 and 2600 A; bromine also exhibits a green fluorescence with an intensity about 1/300th that of the iodine fluorescence; in accordance with this trend,fluorescencein chlorine, only recently reported164 near 5000 A, is weaker still. When iodine vapour is illuminated with appropriate monochromatic light, e.g. Hg 5461 A, it emits afluorescencespectrum which consists, according to the conditions of excitation, of a series of lines (Wood's resonance series) or of bands. A progression of lines results from the overlapping of a single rotational level of a given electronic transition with the exciting line; further series are developed if the irradiating line is broad enough to cover more than one absorption line. The occurrence of fluorescence implies that the molecule remains in the excited state for periods between 10~9 and 10~6 sec; measurements with a Kerr cell give a lifetime of (10 ±0-1) x 10 ~8 sec for the excited state leading to the visible fluorescence of iodine. To maintain an excited state which is sufficiently long-lived to promote fluorescence, 164 w. Hölzer, W. F. Murphy and H. J. Bernstein, /. Chem, Phys. 52 (1970) 399,469. i« M. Suzuki, T.Yokoyama and M.Ito,/. CÄem.PÄ^.50(1969)3392;51 (1969) 1929; J.E.Cahill and G. E. Leroi, ibid. p. 4514.

79Br»lBr 79ßr8iBr

79Br?9Br

79Br8iBr

35C12

Molecule

nltt \

Amlu

a

Ββ (cm-i)

17,672-6 0

13,814 0

153 324-24

167-55

(480) (220) (330)

2-7 1172

1-625 -0015 0-342x10-2

-0010 00657 0081079

0059579

01626 0-24399

UNSTABLE -ΟΟΟ67

We (cm-i)

Vmax^ 23,525 5-3 2-675

(cm-i) STABLE STABLE UNSTABLE STABLE UNSTABLE

259-5 559-72

(cm-i)

(-75,000) (-67,700) "max > 64,100 (-58,000) v max ~30,300

Te (cm-i)

F(U2tlg) (66,500) E(mg) (61,444) D(*nu) (55,535) (47,000) C(3X ltt +) »'max-'44,500 c? Vmax~ 24,300 Amu Vmax~20,740 Bm0+ 0 w ~ 18,630) Am! 15,902-51 Bm0+uc

Amo+u*

3

£(HI,) D(*Ilu)

State

000155 000030405

00004868

000212 0-00149

(cm-i)

205x10-8

309x10-8

2-365xl0-7 -I85XIO-7

De (cm-i)

Numerous continua in the visible and ultraviolet regions

Comments

2-6777 Convergence limit at 19,578 cm-i 2-55 2-2809

>Continua

Numerous diffuse emission bands and continua in the visible and near-ultraviolet

2-43 5 1 Convergence limit at 1-9881 20,880 c m - i

re (Ä)

TABLE 12. SPECTROSCOPIC PROPERTIES OF DIFFERENT STATES OF THE DIATOMIC HALOGEN MOLECULESB

ΖΐΣ, *

Amlu +

Bm0+U*

C(3E ltt +) £ΡΣ,+) />(1Σ„ + )

F(mu)

#(nlf0+«) ΑΠι.ο«) /(n l f 2 u) H(UU2u) C?(^tt+)

M(U2.U)

State

(120) (360) (215) (260) (212) (260) (206) 1651 96-2 (90) 102-2 104 125-53

44-0 214-52

11,888 0

ωβ (cm-i)

(58,620) (55,930) (51,528) (62,780) (59,218) (57,770) (56,937) 51,708 46,488-7 (44;960) 41,407 33,744 15,769-48

Te (cm-i)

1·0ι 0-60738

0-595 0-49 (017) 0-34 0-2 0-7339

ωβχβ

(cm-i)

+0008 -1-307x10-3

000413

-00035

ωβ^β

(cm-i)

0037389

0028873

Be (cm-i)

00001210

00001345

«e

(cm-i)

4-54x10-9

-3-5x10-9

De (cm-i)

2-666

3033

re

(A)

Convergence limit at 20,037 c m - i

Cordes bands

1 Numerous diffuse > emission bands I and continua

Comments

a G. Herzberg, Molecular Spectra and Molecular Structure. I. Spectra of Diatomic Molecules, pp. 512-541, van Nostrand (1950); Supplement to Mellor's Comprehensive Treatise on Inorganic and Theoretical Chemistry, Supplement II, Part I, Longmans, London (1956); for most recent review see ref. 161. b A. E. Douglas, Chr. Kn. Möller and B. P. Stoicheff, Canad. J. Phys. 41 (1963) 1174; M. A. A. Clyne and J. A. Coxon, /. Mol. Spectroscopy, 33 (1970) 381. c J. A. Horsley and R. F. Barrow, Trans. Faraday Soc. 63 (1967) 32; J. A. Coxon, / . Mol. Spectroscopy, 37 (1971) 39. d M. A. A. Clyne and J. A. Coxon, /. Mol. Spectroscopy, 23 (1967) 258; J. A. Horsley, ibid. 11 (1967) 469. β R. J. LeRoy and G. Burns, / . Mol. Spectroscopy, 25 (1968) 77. * J. I. Steinfeld, J. D. Campbell and N. A. Weiss, /. Mol. Spectroscopy, 29 (1969) 204. « D. H. Rank and B. S. Rao, /. Mol. Spectroscopy, 13 (1964) 34. Values in parentheses are uncertain.

127I2

Molecule

Table 12 (cont.)

1184

CHLORINE, BROMINE, IODINE AND ASTATINE: A. J. DOWNS AND C. J. ADAMS

it is a necessary condition that other modes of deactivation of the excited state do not take place in times shorter than 10 ~9 to 10 ~6 sec. Relaxation of this condition, as by the addition of foreign gas or by the application of a magnetic field, leads ultimately to quenching of the fluorescence. The absorption spectra of bromine and iodine vapours in the region 2000-3000 A have recently been found to exhibit significant variations with temperature and concentra­ tion166»167, findings which have been interpreted in terms of the equilibrium 2X2 τ^ Χ4

The temperature-dependence of the equilibrium constant over the range 150-420°C implies that for iodine ΔΗο605°κ is -2-9 ±0-4 kcal mol - 1 and Δ5 ο 605 ο κ is —14-4 ± 2 eu, while at 240°C and 2-5 atm pressure 1 -4 mol%I4 is present in the vapour; PVT data for bromine lead to the values Δ/ί ο 400 ο κ = —2-6 kcal mol - 1 andÄS,°40o0K = ~ 15 eu. The thermodynamic results are consistent with the formation of a complex involving van der Waals' interaction between the X2 partners which are thus able to move relatively freely with respect to each other. That certain features in the ultraviolet-visible spectra of solutions of iodine in carbon tetrachloride, diethyl ether or 1,2-dichloroethane do not obey Beer's Law has likewise been attributed to the production of the I4 aggregate with a formation constant ranging from 2-5 (CCI4) to 0-5 1 mol"1 (Et20)i68-no. The molecular ions X 2 + have been variously identified, as by their emission spectra when the gaseous halogen is excited in a discharge tube; by mass-spectrometric analysis of the behaviour of the parent molecule under the action of electron-impact; and by measure­ ments of the photoelectron spectra. This last method provides what are probably the most reliable and precise estimates of the ionization potentials of the diatomic molecules171; these values are contained in Tables 11 and 39. The dissociation energy of the molecular ion DQ(X2 + ) is related to D0(X2) and the ionization potentials I(X) and /(X2) by A>(X2+) = A)(X2)+/(X)-/(X 2 )

The fact that DQ(X2 + ) so derived is consistently higher than D0(X2) (see Table 39) confirms the anti-bonding character of the 7rg orbital from which ionization takes place. Another notable feature is that the dissociation energy of F 2 + (3-42 eV) is lower than that of Cl2 + (3-99 eV), just as the dissociation energy of F 2 is lower than that of Cl2. Detailed rotational and vibrational analysis172 of the banded spectrum exhibited in emission by the Cl2 + ion has afforded the parameters given in Table 39. Though spectroscopic and magnetic properties of condensed phases containing the Br2 + and I2 + ions have recently been described, significant perturbation of the ground states of the molecular ions are to be expected as a result of ion-ion or ion-solvent interactions. The influence of chemical environment on the formation and behaviour of halogen cations is more appropriately treated elsewhere (Section 4, p. 1340). 16 <> A. A. Passchier, J. D. Christian and N. W. Gregory, / . Phys. Chem. 71 (1967) 937. 167 A . A. Passchier and N. W. Gregory, / . Phys. Chem. 72 (1968) 2697; M. Tamres, W. K. Duerksen and J. M. Goodenow, ibid. p. 966. 168 L. J. Andrews and R. M. Keefer, Adv. Inorg. Chem. Radiochem. 3 (1961) 91. 169 R . s. Mulliken and W. B. Person, Molecular Complexes, p. 141. Wiley, New York (1969). 170 D . D. Eley, F. L. Isack and C. H. Rochester, / . Chem. Soc. (A) (1968) 1651. 171 A. W. Potts and W. C. Price, Trans. Faraday Soc. 67 (1971) 1242; S. Evans and A. F. Orchard, Inorg. Chim.Acta, 5 (1971) 81. 172 F. P. Huberman, / . Mol. Spectroscopy, 20 (1966) 29.

PHYSICAL PROPERTIES OF THE MOLECULAR HALOGENS

1185

The Condensed Phases The molecular lattices adopted by solid chlorine, bromine and iodine are isostructural. In contrast with the allotropy of neighbouring elements in Group VI, there is no good evidence to suggest the existence of allotropes of solid chlorine, bromine or iodine under normal conditions. The unit cell of the crystals is invariably orthorhombic with the space group Cmca containing four molecules and having the dimensions listed in Table 11173. The intramolecular X-X distance in solid chlorine or bromine is not significantly different from that of the gaseous molecule, but refinement of the crystal structure of iodine at 110°K shows that at 2-715 A the I-I bond is appreciably longer than that of the gaseous molecule (2·666 A). Figure 13 depicts the two-dimensional network formed by nearly linear chains

/N/

x X

FIG. 13. The two-dimensional network of iodine molecules in the (100) plane of solid iodine. [Reproduced with permission from F. van Bolhuis, P. B. Koster and T. Migchelsen, Acta Cryst. 23 (1967) 90.]

I — I—I — I, each iodine atom being involved in two nearly perpendicular chains. The angles of approximately 90° suggest that, to afirstapproximation, the bonds in intersecting chains are formed by orthogonal 5p orbitals of the iodine atom. Each chain can then be described in terms of a model involving four-centre six-electron bonds174; alternatively the space group of the unit cell can be exploited to define the matrix integrals for a model in which there is two-dimensional delocalization of /^-electrons, and which accounts semiquantitatively for the nqr properties, crystal energy and conductivity of solid iodine175. In certain respects, however, an even more apposite description is to be found in the band model, which has been successfully applied to solid iodine on the assumption that electronexchange between layers is negligible and that only p functions are used for binding by the iodine atoms176. The relatively short intermolecular Χ···Χ distances in solid chlorine and bromine bespeak situations similar in type, if not in degree. Calculations suggest that the orthorhombic structure adopted by the halogens itself testifies to relatively strong intermolecular forces since cubic or hexagonal structures would otherwise seem to be energetically more favourable177. That the intermolecular forces assume increasing importance in the series 173 j . Donohue and S. H . Goodman, Acta Cryst. 18 (1965) 568; F. van Bolhuis, P. B. Koster and T. Migchelsen, ibid. 23 (1967) 90. 174 R . j . Hach and R. E. Rundle, / . Amer. Chem. Soc. 73 (1951) 4321; G. C. Pimentel, / . Chem. Phys. 19 (1951) 446. 175 J. L. Rosenberg, / . Chem. Phys. 40 (1964) 1707. 176 R. Bersohn, / . Chem. Phys. 36 (1962) 3445. 177 s . C. Nyburg, / . Chem. Phys. 40 (1964) 2493; K. Yamasaki, / . Phys. Soc. Japan, 17 (1962) 1262.

1186

CHLORINE, BROMINE, IODINE AND ASTATINE: A. J. DOWNS AND C. J. ADAMS

Cl2, Br2, I2 is well illustrated by the variation in the ratio of the shortest intermolecular X · · -X to the intramolecular X-X distance (Table 11). Evidence of this same general trend is also to be found in the following features. (i) As the atomic number advances, so the volatility of the solid halogen decreases, while there is an increase in the melting point and enthalpies of fusion and sublimation; likewise the volatility of the liquid is depressed and the boiling point and heat of volatili­ zation are elevated. These variations follow the general pattern defined by van der Waals' interactions, which are functions primarily of polarizability. (ii) The vibrational frequency and force constant of the X-X molecule (X = Cl or Br) are significantly reduced in the transition from the gas to the solid phase (see p. 1181). Combined with the observed vibrational frequencies of the lattice and the crystal-field splitting of the intramolecular stretching vibration, the findings indicate stronger inter­ molecular interaction in solid bromine than in solid chlorine. A similar pattern has also been deduced from studies of the far-infrared spectra of crystalline chlorine, bromine and iodine, whence intermolecular force constants have been calculated1™. (iii) Measurements of the quadrupole resonance spectra of the 35C1, 3<7C1, 79Br and 12?I nuclei in the respective solid halogens afford the results given in Table 11. For solid bromine and iodine, the asymmetry parameter η for nuclear-quadrupole coupling is found to depart substantially from the zero value characteristic of a molecule in a Σ state. This signifies that the electricfieldat the halogen nucleus is not symmetric about the bound axis, a circumstance presumably dictated by intermolecular binding. The fact that the quadrupole coupling constants found for the crystalline halogens are only slightly less than those of the free atoms gives good grounds for the belief that the bond in the X 2 molecule derives from the overlap of p-orbitals with little, if any, admixture of .y-orbitals. More detailed studies give notice that significant orbital interaction contributes to the intermolecular forces176»1<79. (iv) The layer structure of solid iodine is well characterized by the marked anisotropy of the dielectric constant, electrical conductivity and diamagnetic susceptibility (see Table 11), which have been measured using single crystals of iodine. Monocrystalline iodine thus emerges as a two-dimensional semiconductor having a band-gap estimated to be ca. 1 -3 eV18<>. It also possesses unique photoconducting properties shown to be the result of hole- rather than electron-conduction, the dominant hole traps being shallow and lying 0-4-0'5eV above the valence band edge181. Under high pressures solid iodine suffers a remarkable increase in electrical conductivity157»182; measurements carried out on single crystals indicate a changeover to metallic behaviour in a direction perpendicular to the plane of the molecules at pressures in excess of 160 kbar and a corresponding changeover in directions parallel to this plane at pressures in excess of 220 kbar. At pressures between 160 and 220 kbar the iodine crystal is a semiconductor in one direction and a metal in the other, after the fashion of graphite at atmospheric pressure. These pressures are to be compared with the thresholds for metallic 178 s . H . Walmsley and A . Anderson, Mol. Phys. 7 (1963-4) 411. 179 C . H . Townes and B . P . Dailey, / . Chem. Phys. 20 (1952) 3 5 ; N . N a k a m u r a a n d H . Chihara, / . Phys. Soc. Japan, 22 (1967) 201; E. A . C. Lücken, Nuclear Quadrupole Coupling Constants, p. 288. Academic Press, London and N e w Y o r k (1969). 180 N . N . Kuzin, A . A . Semerchan, L . F . Vereshchagin a n d L . N . Drozdova, Doklad. Akad. Nauk S.S.S.R. 147 (1962) 7 8 ; L . F . Vereshchagin a n d E . V. Zubova, Fiz. Tverd. Tela, 2 (1960) 2776. I» 1 A . Many, M . Simhony, S. Z . Weisz a n d J. Levinson, Phys. Chem. Solids, 22 (1961) 285. 182 L . F . Vereshchagin, A . A . Semerchan, S. V. Popova a n d N . N . Kuzin, Doklad. Akad. Nauk S.S.S.R. 145 (1962) 757; B . M . Riggleman and H . G . Drickamer, / . Chem. Phys. 38 (1963) 2721.

PHYSICAL PROPERTIES OF THE MOLECULAR HALOGENS

1187

behaviour of 130 and 185 kbar in selenium and stannic iodide, respectively. The develop­ ment of metallic properties is confirmed by the positive temperature coefficient of the electrical resistance of iodine at high pressures183. The effect of such high pressures on the lattice parameters of solid bromine and iodine184 has also been investigated. Whereas bromine crystals remain orthorhombic, the dimensions of the iodine lattice imply a collapse of the normal structure to a more close-packed one with an increased number of near neighbours surrounding a given iodine atom. Mechanisms for the approach to this condition in and perpendicular to the ac plane have been outlined184; whether the I 2 molecules retain their identity and conduction is a result of band overlap or whether the molecules dissociate to give an aggregate of atoms possessing an unfilled conduction band has yet to be resolved. With iodine, therefore, the increased number of relatively low-lying atomic orbitals and the reduced effective nuclear charge experienced by the valence electrons lead to an obscuring of the normal distinction between a molecular and a continuous solid structure. In this respect iodine shows a distinct affinity to elements like selenium and phosphorus, an affinity which is furthered by the relatively high electrical conductivities or semiconductor properties which may be induced in iodine by certain donor moieties, e.g. I3~ 1 8 5 , N,N'-diphenyl-/?phenylenediamine186, perylene and pyrene43, though the precise mechanism of conduction may be open to question in at least some of these systems. Such donor-acceptor interaction encourages the formation of cation radicals representing a source of "holes" and of anion radicals representing a source of electrons. On this basis the electron conductivity of an aggregate is enhanced by a high concentration of the acceptor, viz. iodine, or by incorporation in a polyacceptor network of the type encountered in the black starch-iodine complex186. The band gap appears to depend on the separation between the excited charge-transfer and ground states of the donor-acceptor system; in common with the activation energy for conduction, this separation is less than 0-5 eV for some solid complexes of bromine and iodine. Under normal pressures the liquid phases of chlorine, bromine and iodine each span a temperature range of about 70°C. According to the data of Table 11, liquid chlorine and bromine are both characterized by meagre specific conductivities and low dielectric constants; by contrast, liquid iodine is a much better conductor with a significantly larger dielectric constant. The properties of the liquids suggest that liquid chlorine and bromine are relatively poor solvents, whereas iodine may be expected to function more efficiently in this respect. Extensive studies32»54»187 have confirmed the validity of this generalization. Thus, whereas alkali-metal halides are, for the most part, sparingly soluble in liquid chlorine or bromine, the iodides of the alkali metals, but not those of the alkaline-earth metals, are readily soluble in liquid iodine. Bromides of large cations, e.g. caesium and tetrabutylammonium, dissolve in bromine to give highly viscous solutions, the equivalent conductance of which increases with increasing concentration; the solubility of caesium bromide provides an efficient process for its separation from the other alkali-metal bromides54. On the basis of conductivity measurements on iodides in liquid iodine, the self-ionization /ll 2 ^ Ι + + Ι 2 η - Γ 183 B. M. Riggleman and H . G . Drickamer, / . Chem. Phys. 37 (1962) 4 4 6 . 184 C. E. Weir, G. J. Piermarini and S. Block, / . Chem. Phys. 5 0 (1969) 2089; R. W. Lynch and H . G. Drickamer, / . Chem. Phys. 45 (1966) 1020. 185 T. Sano, K. O k a m o t o , S. Kusabayashi and H . Mikawa, Bull. Chem. Soc. Japan, 4 2 (1969) 2505. 186 V. Hädek, / . Chem. Phys. 4 9 (1968) 5202. 187 A . G. Sharpe, Non-aqueous Solvent Systems (ed. T. C. Waddington), p. 285, Academic Press (1965).

1188

CHLORINE, BROMINE, IODINE AND ASTATINE: A. J. DOWNS AND C. J. ADAMS

has been postulated188, the conduction mechanism probably depending primarily on the facile transfer of iodide ions between iodine molecules. According to the proposed scheme for self-ionization, alkali-metal iodides, which readily form polyiodide anions, can be regarded as bases and iodine monohalides as acids. Conductimetric titration of, for example, KI with IBr in molten iodine shows a sharp break at a 1:1 molar ratio, which would corre­ spond to the process I 3 -+IBr->Br-+2I 2

Such a reaction can also be followed potentiometrically. Outside the vapour phase, however, the simple I + species has yet to be characterized; on present evidence all that can be said is that the simplest, plausible basis for the self-ionization of iodine appears to involve the formation of I3 + and I 3 ~ ions, but that more complex species may well participate. Solvolytic reactions in liquid iodine, e.g. KCN+I 2 ^KI+ICN and amphoteric behaviour, e.g. HgI2+2KI^K2HgI4 have also been described. Many molecular halides which do not otherwise react with chlorine, bromine or iodine are soluble in the liquids presumably as the result of relatively strong interactions between the solute, functioning mainly as the donor partner, and the solvent molecules primarily acting in the role of acceptor. Since most studies of such donor-acceptor interactions relate to the solid phase or to conditions in which the donor molecule is present in large excess, detailed discussion of this subject is more aptly taken up in the next subsection. The addition compounds formed by bromine with organic bases such as pyridine, quinoline, acetamide and benzamide form conducting solutions in bromine, probably giving (base)nBr + and Br3 ~ ions, but the conductivity varies with time, suggesting that bromination of the organic molecule is taking place and at the same time confusing the interpretation of the data187. Sulphur, selenium and tellurium are said187 to dissolve in liquid iodine without chemical change. 2.7. CHEMICAL PROPERTIES OF THE HALOGENS Redox Properties: Aqueous Chemistry A considerable degree of order can be found in the reactions of chlorine, bromine and iodine in aqueous solution if full and proper use is made of the standard oxidation potentials appropriate to the elements and their ions. Direct measurements of standard potentials have been made for the couples |C12/C1 ~, £Br2/Br~ and JI 2 /I ~ with the results given in Table 11, which signify a marked decrease in the oxidizing power of the free halogen in the series F 2 > Cl2 > Br2 > I 2 . There are very few measured potential data for systems involving halogens in other oxidation states, and nearly all the information available159»189 is derived from thermochemical measurements and estimates or from purely chemical evidence. The doubtful reliability of some of the thermal data is discussed in a paper on the heat of formation and the entropy of the bromate ion in solution190. 188 D . J. Bearcroft and N . H . Nachtrieb, / . Phys. Chem. 71 (1967) 316. 189 W. M. Latimer, The Oxidation States of the Elements and their Potentials in Aqueous 2nd edn., Prentice-Hall, N e w York (1952). 190 H . C. Mel, W. L. Jolly and W. M. Latimer, / . Amer. Chem. Soc. 75 (1953) 3827.

Solutions,

CHEMICAL PROPERTIES OF THE HALOGENS

1189

Most reactions involving oxyhalogen anions also involve hydrogen or hydroxide ions and the position of an equilibrium such as the disproportionation of bromine 3Br 2 +3H 2 0 ^ 5Br - + Br0 3 " + 6H +

is very dependent upon the pH of the medium. The potential diagrams originally given by Larimer*^ for the halogens at unit hydrogen ion or at unit hydroxide ion activities are reproduced in a modified form in Fig. 14 (see also Fig. 2), giving what are believed to be the most reliable guides to the aqueous chemistry of the elements at present available. +7

+5 Acid solution a H + = 1

clo4_i±^

+3

Oxidation state -}-l + 1-46

0

-1

Jo- - ί ± 2 ^ Η α ο 2 - ± ^ Η ο α - ± t £ ~ Ji2 +136 -q-

I

t

ciq, +1-27 z

+1-15

+1-51 B r o 4 -i±Z^

±1*

BK>J

^ ΗΟΒΓ ± ! ^ _ £ -±±5Z_ B r 4-1-20

' TT T ^ ca. +l·? ^ H5IO^—' »- I 0 3

+114

____ *- HOI

1

+1-45 ^ ' »· I2(c)

+ 0-50

+0-54, T . *- I

+ 0· 89

co-+M2* do,- +™+ co- -±£&- i,o- + 2 ^ * a 2

-±Ι^_ΟΓ

-CIOH

500 ^ ^ + W6 +061 •+0-93. Br04-^^i*

Λ BiOj

±^54 =

+045 „ ΒΓθ_ ► BrO- - i - ^ - ~ Br2

+107 ~"' "'

.

1_ ft

Μ)·76

+0-29

Η&-!±±*1

ΐό5

±^

.. κ>- _i±45_ I

. Hc)

ί 0 49

+ ^ j

FIG. 14. The standard potentials for systems involving chlorine, bromine and iodine and their ions in aqueous solution, E° in volts.

In the sequence of changes £X2(standard) -> X(g) -> X-(g) -* X"(aq)

the values of Δ#/°[Χ -(g)], and consequently of AGy°[X-(g)], do not vary widely from halogen to halogen. Accordingly the greatest part of the differences between the redox

1190

CHLORINE, BROMINE, IODINE AND ASTATINE: A. J. DOWNS AND C. J. ADAMS

properties of the couples £X 2 /X" arises from the substantial variations in the hydration energies of the anions which follow the sequence Cl~ > Br~ > I - , diminishing as the size of the ion increases. In liquid ammonia the following standard potentials (at 298°K) have been recorded^: £C12/C1-, +1-91 V; £Br2/Br-, +1-73 V; flfe/I", +1-26V. It is of interest that the increase in the potential in moving from water to liquid ammonia is 0-1 V larger for iodine than for the other halogens. This suggests that, relative to water, the iodide ion coordinates more strongly with liquid ammonia than do the other anions, a conclusion that is in accord with the noticeable solubility of iodides in the latter solvent. The couple £X 2 /X" i s reversible for all three halogens, which are characteristically fairly rapid in their oxidizing action, very much faster than molecular oxygen for example. Because of the mechanism of formation of the X2 molecule by oxidation of the X - anion, significant activation energies and hence overvoltages must exist; measurements are said to indicate the following order of overvoltages at a platinum electrode: I2 > Cl2 > Br2. Were it not for the appreciably higher oxygen overvoltage, chlorine would not be evolved in the electrolysis of aqueous chloride solutions at low current densities; accordingly chlorine overvoltages must necessarily be measured at an electrode which is polarized with respect to oxygen. This phenomenon is of great technical importance in relation to the electro­ lytic production of chlorine. The mechanism of oxidation is doubtless complex, possibly involving the formation of hypohalite surface compounds which react with halide ions to give the molecular halogen. The relative proportions of oxygen and chlorine evolved in the electrolysis of chloride solutions are known to be influenced by the presence of metal ions, e.g. Mn 2+ , which preferentially catalyse one or other of the discharge reactions. The halogens are all to some extent soluble in water. In acid, neutral or alkaline solution the standard potential of the couple \02, 2H + /H 2 0 is +1 -23, 0-81 or 0-40 V, respectively. Depending on the precise conditions of pH, all three halogens should therefore be capable of oxidizing water. In contrast with the behaviour offluorine,the reaction tends to be slow, however, and disproportionation is the initial result: X2+H20 ^HOX+H + +XThe weak hypohalous acid so produced then undergoes slow decomposition: e.g. HOC1 -*HC1+K>2 or

2HOBr -> B r 2 + H 2 0 + £ 0 2

so that oxidation of water is the overall result. However, in the case of bromine and iodine, creation of HOX and HX by disproportionation raises the acidity and eventually stops the reaction for, as the potentials show, the oxidation of water by bromine or iodine does not proceed at unit activity of hydrogen ions. For saturated solutions of the halogens in water at 298°K the compositions are as shown in Table 13. There is an appreciable concentration of hypochlorous acid in a saturated aqueous solution of chlorine, a smaller concentration of hypobromous acid in a saturated solution of bromine, but only a very meagre concentration of hypoiodous acid in a saturated solution of iodine. It is evident that, because of the unfavourable equilibria, the reaction of the halogen with water does not constitute a suitable method for preparing aqueous solutions of the hypohalous acid. Nevertheless, the yield of hypohalous acid may be improved by judicious choice of conditions whereby the pH is increased, for example by 191 W. L. Jolly, / . Chem. Educ. 33 (1956) 512.

CHEMICAL PROPERTIES OF THE HALOGENS

1191

TABLE 13. EQUILIBRIUM CONCENTRATIONS IN AQUEOUS SOLUTIONS OF THE HALOGENS AT 298°Ka

Cl2

Property 1

Total solubility (mol Γ ) K\ = [X2(aq)]/[X2(standard)] K2 = [H+][X-][HOX]/[X2(aq)] (moP 1-2) [X2(aq)](moll-i) [H+] = [X-] = [HOX](mol l"i) Thermodynamic properties of X2(aq): AG/°(kcalmol-i) AÄ>°(kcalmol-i) 5°(caldeg-imol-i)

b

00921 0062 4-2x10-4 0062 0030 + 1-65 -5-6 29

Br2

I2

0-2141 0-21 7-2x10-9 0-21 115x10-3

00013 00013 20x10-13 00013 6-4x10-6

+0-94 -0-62 31-2

+ 3-92 + 5-4 32-8

ft

Ref. 32; F. A. Cotton and G. Wilkinson, Advanced Inorganic Chemistry, 2nd edn., p. 569, Interscience (1966); National Bureau of Standards Technical Note 270-3, January 1968. b Cl 2 gas at 1 atm pressure.

the use of a suspension of mercuric oxide: 2 X 2 + 2 H g O + H 2 0 -> HgO,HgX 2 +2HOX

All of the hypohalous acids are rather unstable, being also relatively strong oxidizing agents' especially in acid solution. The analogous sulphur compounds HSX appear to be even less stable on the very limited evidence that is available. Nevertheless, bromine is said192 to react under rigidly controlled conditions with hydrogen sulphide in chloroform or dichloromethane solution according to the scheme: Br 2 +H 2 S -> HBr+HSBr

and the isolation of the salt NH4SBr at low temperature has been reported. In alkaline solution the halogens are converted to the corresponding hypohalite ions in accordance with the general reaction X 2 +20H- ^ X

+XO+H20

The equilibrium constant for this reaction is invariably favourable—7-5 x 1015 for chlorine, 2 x 108 for bromine and 30 for iodine—and the reaction is rapid. Thus the hydrolysis of chlorine is practically complete in alkaline solution, but the balance is reversed in acid solution. The preparation of bleaching powder and sodium hypochlorite and their action as bleaching agents provide an important illustration of the operation of this equilibrium. However, the situation is complicated by the tendency of the hypohalite ions to dispropor­ tionate further in basic solution to produce the corresponding halate ions: 3XO- ^ 2 X + X 0 3 -

For this reaction the equilibrium constant is in each case very favourable, viz. 1027 for ClO ~, 1015 for BrO~ and 102<> for IO -. Thus the actual products obtained when a halogen is dissolved in alkaline solution depend on the rates at which the hypohalite ions initially produced undergo disproportionation. The disproportionation of ClO ~ is slow at and below room temperature, but becomes fairly rapid at ca. 75°C. By the choice of appro­ priate conditions good yields of the chlorate ion can be secured, as in the commercial !92 M. Schmidt and I. Löwe, Angew. Chem. 72 (1960) 79.

1192

CHLORINE, BROMINE, IODINE AND ASTATINE: A. J. DOWNS AND C. J. ADAMS

production of chlorates. The disproportionation of BrO ~ is moderately fast even at room temperature, and at temperatures of 50-80°C the Br0 3 ~ ion is formed quantitatively according to the equation 3Br 2 +60H" -> 5Br" + B r 0 3 " + 3H 2 0

Thus bromine is thermodynamically stable in acid solution with respect to disproportiona­ tion into bromate and bromide, but unstable at a pH of 10; advantage has been taken of this behaviour in an important stage of one process for the extraction of bromine from sea water (see p. 1138). The rate of disproportionation of IO ~ is so fast at normal temperatures that the ion has so far eluded proper characterization. With strong base, therefore, iodine reacts quantitatively to give iodate and iodide ions. Halite ions and halous acids do not feature in the hydrolysis of the halogens. The disproportionation 4C10 3 - ^ C 1 - + 3 C 1 0 4 2

has an equilibrium constant of 10 <>, but takes place only very slowly in solution even near 100°C. Such kinetic barriers are an outstanding characteristic of the formation and chemical properties of the perchlorate ion. Likewise, although perbromate emerges as a more powerful oxidant than either perchlorate or periodate, it appears to be sluggish in formation and reaction19*. Unlike chlorate, iodate is stable with respect to reactions such as 7 I 0 3 - + 9 H 2 0 + 7H + ^ I 2 + 5 H 5 I 0 6 , K = 10 "85 4 I 0 3 - + 3 0 H - + 3 H 2 0 ^ I" +3Η 3 Ι0 6 2 ", K = 10"44

so that, irrespective of rate, disproportionation with the formation of periodate is not favoured. Many other redox reactions of the molecular halogens are readily comprehensible on the basis of the thermodynamic data of Fig. 14; some typical reactions are listed in Table 14. Thus the decrease in oxidizing power with increasing atomic number leads to replacement reactions such as Cl 2 +2Br" -+Br 2 +2C1-

in which a halogen of lower atomic number displaces one of higher atomic number. Likewise redox couples with values of E° less than that of the couple £X 2 /X" are usually subject to oxidation by the halogen X2, though kinetic factors may supervene in some cases. For example, in aqueous solution chlorine, bromine or iodine oxidizes, inter alia, nitrite to nitrate, arsenite to arsenate, sulphite to sulphate, hydrogen sulphide to sulphur, Sn2 +(aq) to Sn4 +(aq) and [Fe(CN)ö]4~ to [Fe(CN) 6 P - . However, although the balance of the equilibrium 2Fe2 + (aq)+X 2 ^ 2Fe3 + (aq)+2X "

favours the ferric state when X = Cl or Br, in neutral or acid solution ferric ions oxidize iodide to iodine. Again, whereas thiosulphate ions are oxidized to sulphate by chlorine or bromine, the familiar and analytically important reaction 2S 2 0 3 2~ + I 2 -> S 4 0 6 2" + 2 1 -

takes place with iodine. In fact the standard redox potential of the couple SO42 "/iS2C>32 ~ (ca. +0-3 V) is such that iodine, in common with the other halogens, should be able to 193 E . H. Appelman, Inorg. Chem. 8 (1969) 223.

Selenite ions

Thiosulphate ions

Sulphite ions

Hydrogen sulphide

Arsenite

Hypophosphite, phosphite and hypophosphate ions

Azide ions

Hydrazine

Ammonia

Nitrite ions

Halide ions

Reagent

Se0 3 2- +C12 + H 2 0 -> Se0 4 2- +2H + +2C1"

X 2 +2Y- ->2X"+Y 2 (X = Cl, Y = BrorI;X = Br, Y = I) The reaction may proceed further, particularly under alkaline conditions: e.g. I 2 +5C1 2 +6H 2 0 -> 2I0 3 " + 10C1" + 12H+ N 0 2 - + i X 2 + H 2 0 -> NO3- + X - +2H + (X = Cl, Br or I) 2NH 3 +3X 2 ->N 2 +6H + +6X(X = ClorBr) N 2 H 4 +2X 2 -> N 2 +4H + +4X(X = Cl, Br or I) 2N 3 -+X 2 ->3N 2 +2X(X = Br or I) Typically: X 2 +H 2 P0 3 - + H 2 0 -> Η 2 Ρ0 4 - +2Η + +2Χ" (X = Cl, Br or I) H 3 As0 3 +X 2 +H 2 0 ->H 3 As0 4 +2H + +2X~ (X = Cl, Br or I) H 2 S+X 2 ->S+2H + +2X" (X = Cl, Br or I) S0 3 2- + X 2 + H 2 0 -> SO42- +2H + +2X(X = Cl, Br or I) S2032" +4X 2 +5H 2 0 -> 2HS0 4 " +8H + +8X(X = Cl or Br) 2S 2 0 3 2-+I 2 ->S 4 0 6 2 -+2I-

Reaction

This reaction, normally carried out in acid solution in the presence of iodide ions, is of particular importance in the volumetric estimation of iodine. Neither bromine nor iodine is a sufficiently strong oxidizing agent under normal conditions to effect this change.

However, concentrated H 2 S0 4 oxidizes Br~ or I" to the free halogen.

The reaction with iodine has been used to estimate phosphites and hypophosphites. The reaction between I 2 and arsenite, which is quantitative at pH 4-9, is an important analytical reaction. The sulphur may be further oxidized to S0 4 2 _ by chlorine or bromine.

The reaction is catalysed by traces of SCN - , S 2 0 3 2-, S2~, thioketones or mercaptans.

Iodine undergoes a complicated reaction, the course of which depends upon the conditions; products include NH4I and NI3.

Bromate and iodate ions are formed.

Such reactions are important in the extraction of bromine and iodine and in the iodometric determination of free chlorine and bromine.

Comments

TABLE 14. SOME REDOX REACTIONS OF THE MOLECULAR HALOGENS IN AQUEOUS SOLUTION*

b

β

I 2 +5I0 4 - +H 2 0 -» 7I0 3 - +2H+

Cl2+I03"+2Na + + 30H- ->Na2H3I06 1 +2Q"

3I2+5C103-+3H20-> 6IO3-+6H++5C12ΐ 2 +ιο 3 -+ιοα-+6Η + -^5ia 2 -+3H 2 o

ci 2 +2cio 2 - ->2ci-+2cio 2

Refs. 32 and 33. H. Taube and H. Dodgen, /. Amer. Chem. Soc. 71 (1949) 3330.

Perhalate

Halate

Halite

A representative reaction is: X 2 +5C10- + H 2 0 -* 2XO3- + 5C1" +2H + (X = Br or I)

Typically I 2 +8H + + 10NO3- -»2IO 3 - + 10NO2+4H2O

Concentrated nitric acid, persulphate or permanganate Oxyhalogen species: Hypohalite

Ιτοη(Π)

Τίη(Π)

Examples include: HC0 2 - +X 2 -> C 0 2 + H + +2XC 2 0 4 2 - +X 2 -> 2C0 2 +2X~ (X = Br or I) Sn2 + +X 2 ->Sn4 + +2X" (X = Cl, Br or I) 2Fe2+(aq)+X2 ^2Fe3 + (aq)+2X" 2[FeiCN)6]4- +X2 ^ 2[Fe(CN)6p- +2X" (X = Cl, Br or I)

Reaction

Carboxylate ions with reducing properties

Reagent

Table 14 (cont.)

Occurring m strongly acid media, this forms the basis of the Andrews procedure in which potassium iodate is used as an analytical oxidiz­ ing agent. A convenient synthesis of sparingly soluble periodate derivatives; K42O9 and Ag4209 can be prepared in a similar manner.

This represents an important commercial method of producing C102. A tracer study shows that most of the Cl atoms in the C102 are derived from the C102"; an unsymmetrical intermediate CI-CIO2 or Cl-O-Cl-O may be formed.b

Preparative reaction for iodic acid.

Fe(III) favoured by X = Cl or Br; Fe(II) favoured by X = I. In neutral solution I 2 oxidizes [Fe(CN)6]4" to [Fe(CN)6]3", while in strongly acid solution the reverse reaction occurs; this behaviour can be exploited for the estimation both of ferrocyanides and of ferricyanides.

Comments

CHEMICAL PROPERTIES OF THE HALOGENS

1195

effect oxidation to the sulphate ion; the fact that it oxidizes thiosulphate quantitatively to tetrathionate therefore implies that this reaction is very much faster than the oxidation to sulphate. It is also to be noted that some redox reactions which take place in aqueous solu­ tion do not occur in non-aqueous media. For example, iodine and sulphur dioxide do not react when dissolved together in a mixture of anhydrous methanol and pyridine; only when water is added does oxidation of the sulphur dioxide by the iodine take place (see Table 14). This is the basis of the Karl-Fischer reagent used for the determination of small amounts of water 1 ^. The oxidation potentials characteristic of oxyhalogen systems show them without excep­ tion to be strong oxidizing agents, and the oxidation of the molecular halogens is correspond­ ingly difficult. Nevertheless, although the direct conversion of molecular chlorine to individual oxychlorine species can be achieved with but few reagents, the relative stability of the iodate ion is underlined by the facility with which it is formed from iodine by the action of such oxidizing agents as concentrated nitric acid, chlorate, bromate, chlorine, persulphate or permanganate. It seems probable that the diversity of behaviour exhibited by the halogens with respect to such oxidation depends, at least in part, on kinetic rather than thermodynamic barriers. Photolysis of aqueous solutions of the halogens induces reactions that are otherwise slow under normal conditions. As early as 1785 Berthollet recorded the action of sunlight on chlorine water195, while Balard's reports of the properties of bromine water published in 18264o likewise alluded to the action of sunlight. The consensus of these and subsequent studies is that photolysis favours oxidation of the water by the halogen to produce oxygen, X~ and XO3- ions: e.g. 2C1 2 +2H 2 0 -> 4H + +4C1- + 0 2 5C1 2 +5H 2 0 -* 10H + +CIO3- +9C1" + 0 2

Possible mechanisms proposed32 for these reactions hinge on initiation through the step X 2 4-/j v -> 2X·, and on propagation via free radicals such as -OH, Ό 2 Η and ClO. The decomposition of hypohalous acids and of hypohalites is also greatly accelerated by photo­ lysis, the relative contributions of the two modes of decomposition, 2XO- - * 2 X + 0 2 3XO- - * 2 X - + X 0 3 -

being influenced considerably by temperature, concentration, pH, added salts and exposure to the atmosphere. Such catalytic effects stress once again the importance of kinetic and mechanistic factors in the aqueous chemistry of the halogens. In fact, nearly all the positive oxidation states would be denied existence in aqueous solution but for the slowness of decomposition into the molecular halogen (or halide ion) and oxygen. We are still far from having a clear picture of the mechanisms of many reactions of the halogens in aqueous media. However, as a general rule, the acceptor character of the halogen molecule with respect to nucleophilic reagents like H 2 0, OH ~ or X - leads to labile addition compounds, which may well function as essential reaction intermediates. The formation and characteristics of such compounds provide the theme of the next subsection. 194 A. I. Vogel, A Text-book of Quantitative Inorganic Analysis, 3rd edn., p. 944. Longmans, London (1961).

195 c . L. Berthollet, Mem. Acad. (1785) 276.

1196

CHLORINE, BROMINE, IODINE AND ASTATINE: A. J. DOWNS AND C. J. ADAMS

Acceptor Functions: Charge-Transfer Interactions 32 » 33 » 168 » 196-204

In common with molecules like sulphur dioxide, oxygen, the hydrogen halides, quinones and polynitroaromatic systems, halogen molecules of the types X 2 and XY exhibit a primary function as electron-acceptors forming complexes with a wide range of donor species. Not only is this capacity of the halogens at the root of their behaviour as solutes and, in many circumstances, as chemical reagents, its most direct expression, that is in complex-formation, has influenced profoundly the development of our understanding, in particular, of so-called charge-transfer interactions and, in general, of donor-acceptor functions. It is the vacant anti-bonding au orbital of the halogen molecule (see Fig. 11), inevitably more diffuse than the bonding ag orbital, which furnishes the acceptor capacity. The halogen molecules are accordingly to be classified as σ-acceptors; further, because the acceptor orbital is antibonding in type, with the result that its occupancy must cause a weakening of the X-X or X-Y bond, the molecules are, in Mulliken's language204, "sacrificial" in action, prone to form relatively weakly bonded complexes. On the basis of semi-empirical considerations, the following vertical electron affinities (in eV) have been deduced for the halogen molecules205: Cl2, 1*3 ±0-4; Br2, 1-2 + 0-5; I 2 , 1-7 ±0-5. Although the acceptor power of iodine thus appears to be superior to those of bromine and chlorine, variations of electron affinity are almost certainly outweighed in normal chemical situations by other factors arising, for example, from charge-transfer (see below), dispersion and other environmental interactions. In addition to the primary σ-acceptor function of a halogen molecule, the occupied ng orbitals, which afford a π-donor system, exercise a significant secondary influence. This amphoteric behaviour is consistent with the "soft acid" or "class b " character of the halo­ gens in their interactions with species having primarily a donor function; it is also manifest in the relatively strong intermolecular forces in solid iodine and in the aggregates X4 (X = Br or I) reported to exist in the gas phase and in solution166 _ 1 7 °. The list of compounds which form recognizable complexes with the halogens encompasses two types of donor: (i) σ-donors, which possess formally non-bonding electrons. These include a wide variety of nitrogen bases (e.g. aliphatic amines, pyridine and its derivatives, nitriles), oxygen bases (e.g. alcohols, ethers, carbonyl compounds), organic sulphides and selenides (e.g. 1,4-dithian, 1,4-diselenacyclohexane), and certain halide derivatives (e.g. alkyl halides and halide ions). (ii) π-donors, wherein the donor function is performed by bonding 7r-orbitals. These include aromatic systems ranging from benzene to polycyclic hydrocarbons like perylene, as well as molecules containing more localized ττ-charge clouds, e.g. alkenes. In Mulliken's terminology the σ-donors are "increvalent" by virtue of their capacity to donate lone-pair electrons, whereas the ττ-donors are "sacrificial", donation being from a 196 R . s . Mulliken and W. B. Person, Ann. Rev. Phys. Chem. 13 (1962) 107. 197

G. Briegleb, Elektronen-Donator-Acceptor-Komplexe, Springer-Verlag (1961).

198 O. Hassel and Chr. R o m m i n g , Quart. Rev. Chem. Soc. 16 (1962) 1. I " J. N . Murrell, Quart. Rev. Chem. Soc. 15 (1961) 191. 200 s . F. Mason, Quart. Rev. Chem. Soc. 15 (1961) 353. 2 °i J. Rose, Molecular Complexes, Pergamon, Oxford (1967). 202 H . A . Bent, Chem. Rev. 68 (1968) 587. 203 c . K. Prout and J. D . Wright, Angew. Chem., Internat. Edn. 7 (1968) 659; C . K . Prout and B.Kamenar, Molecular Complexes (ed. R. Foster), Paul Elek (Scientific Books) (1973). 204 R . s . Mulliken and W. B. Person, Molecular Complexes, Wiley, N e w York (1969)* 205 w . B. Person, / . Chem. Phys. 38 (1963) 109.

CHEMICAL PROPERTIES OF THE HALOGENS

1197

bonding orbital, with the result that bonding within the donor is weakened by complexformation. Some systems, e.g. pyridine, are functionally capable of acting both as σ- and π-donors206, while the presence of energetically accessible vacant orbitals, e.g. d- or antibonding 7r-orbitals, may impart amphoteric properties, the donor function being supple­ mented by a secondary acceptor role. Outside the limits defined by these categories, molecules like cyclohexane and n-heptane possessing only low-energy occupied σ-orbitals may yet experience with the halogen molecules short-lived interactions so weak as to preclude the identification of a distinct complex but sufficiently strong to produce marked pertur­ bation of the electronic spectra of the components2**4. The equilibrium X2 + D ^ D , X 2

involving the interaction of a halogen and donor species D is rapidly established. Evidence concerning the existence, stability and structure of complex species such as D,X2 has been derived by a variety of physicochemical methods, to be outlined in the following pages. 1. Historical considerations168

It has been known for many years that the colour of iodine solutions varies with the nature of the solvent. Thus, the halogen is violet in media such as the aliphatic hydrocarbons and carbon tetrachloride; in other solvents, including the alcohols, ethers and benzene, it is brown or reddish-brown207. The visible absorption maximum of the violet solutions is located in the 520-540 ηΐμ, region, the overall spectrum being similar to that of iodine in the vapour state. The maximum for brown solutions occurs at shorter wavelengths (460-480 τημ). The solubility of iodine, its heat of solution and the chemical reactivity are generally greater for solvents that give brown solutions than for solvents that give violet solutions. Further, violet solutions often turn brown on the addition of a solvent such as alcohol, whereas brown solutions often turn violet on heating (and brown again on cooling). Although the molecular weight of iodine in both types of solvent corresponds to the diatomic unit I2, an abnormally small freezing-point depression is observed when a small amount of a solvent forming a brown solution is added to a violet solution. Several ideas have been invoked to explain the varying colours of iodine solutions: it has been suggested that the brown solutions contain colloidal particles, solvated molecules or associated molecules. A theory of solvent-solute interaction in terms of the cage theory of solutions has also been put forward208. However, the formation of molecular complexes, suggested as early as 1930209, was first demonstrated rather conclusively by the notable studies of Benesi and Hildebrand210 on the visible and ultraviolet spectra of solutions of iodine in benzene and other aromatic solvents. Most significant was the discovery of an intense ultraviolet absorption band near 300 imx attributable neither to the solvent nor to iodine. By following the changes in the intensity of absorption with concentration, using an inert solvent, it was shown that the new band arises from a 1:1 complex. Subsequent 206 R . s. Mulliken, / . Amer. Chem. Soc. 91 (1969) 1237; I. D. Eubanks and J. J. Lagowski, ibid. 88 (1966) 2425. 207 A. Lachman, / . Amer. Chem. Soc. 25 (1903) 50; J. H. Hildebrand and R. L. Scott, The Solubility of Non-electrolytes, 3rd edn., Reinhold, New York (1950). 208 N . S. Bayliss and A. L. G. Rees, / . Chem. Phys. 8 (1940) 377; A. L. G. Rees, ibid. p. 429; N . S. Bayliss, Nature, 163 (1949) 764; / . Chem. Phys. 18 (1950) 292. 209 M. Chatelet, Compt. rend. 190 (1930) 927. 2io H. A. Benesi and J. H. Hildebrand, / . Amer. Chem. Soc. 70 (1948) 2832; 71 (1949) 2703.

1198

CHLORINE, BROMINE, IODINE AND ASTATINE: A. J. DOWNS AND C. J. ADAMS

studies168 have disclosed similar behaviour in aromatic solvents on the part of chlorine and bromine as well as interhalogens such as iodine monochloride. Because of the relatively feeble interaction between the components, complex-formation can often be detected only by studying the physical properties of solutions in which the complexes are in equilibrium with their components. Nevertheless, the persistence in the vapour phase of certain complexes, e.g. I2,OEt2, l2,SEt2, I2,C6H6 and C5H5N,2I2, is suggested by the ultraviolet211 and mass212 spectra. Further, numerous crystalline adducts of definite stoichiometry have been isolated; examples include Me3N,I2, 4-picoline,I2, hexamethylenetetramine,2Br2, l,4-dioxan,X2(X = Cl, Br or I), acetone,Br2, l,4-dithian,2I2, l,4-diselenacyclohexane,2I2, C6H6,X2 (X = Cl or Br) and [p-Me2N-C6H4]2C = CH2,2X2 (X = Br or I). The structures of some of these solid complexes have been determined by X-ray crystallographic analysis carried out by Hassel and others198, beginning in 1954 with the adduct l,4-dioxan,Br2. The results of such analysis are summarized in a subsequent section (see pp. 1201-6 and Table 16). 2. Physicochemical methods of investigation168»197»201'204 The most distinctive characteristic associated with the formation of a halogen complex is the appearance of a new intense and broad absorption band in the visible or ultraviolet spectrum, generally accompanied by perturbations of the spectral bands arising from electronic transitions of the component molecules. Several methods have been devised for evaluating the formation constant of such a complex from spectrophotometric measure­ ments of optical density as a function of concentration. Determination of the formation constant at more than one temperature affords values for the enthalpy and entropy changes accompanying complex-formation. Listed in Table 15 are formation constants and enthalpies for some representative complexes of iodine. Formation constants are thus seen to span the range 10 _1 to 104 1 mol - 1 , while enthalpies lie in the range 1-13 kcal mol - 1 . The new absorption band characteristic of each complex, the so-called "charge-transfer" band, has a molar extinction coefficient of the order 104 and a half-width typically of 6000 cm - 1 . The interaction of a donor D with a halogen or interhalogen XY implies a certain amount of charge-transfer in the ground state of the complex. Accordingly the vibrational properties of the complex should reflect not only new motions in which D vibrates against XY, but also changes in the strengths of bonds within the D and XY units. Complexformation is thus attended by the following changes in the infrared and Raman spectra of the component species213. (a) New spectral features attributable to vibrations of the D-XY unit may be observed. Thus for Me3N,I2 such a band, representing in large part the N-I stretching vibration, has been located at 145 c m - 1 in infrared absorption. (b) The frequencies and intensities of vibrational bands associated with internal motions of the D and XY molecules may suffer significant changes. For example, with respect to the 2ii F. T. Lang and R. L. Strong, / . Amer. Chem. Soc. 87 (1965) 2345; M. Tamres and J. M. Goodenow, / . Phys. Chem. 71 (1967) 1982. 212 R . C a h a y a n d J. E . Collin, Nature, 211 (1966) 1175. 213 R . F . L a k e a n d H . W . T h o m p s o n , Proc. Roy. Soc. A297 (1967) 4 4 0 ; P . K l a b o e , / . Amer. Chem. Soc. 89 (1967) 3667; J. G e r b i e r a n d V. Lorenzelli, Spectrochim. Acta, 23A (1967) 1469; F . W a t a r i , ibid. p . 1917; Y . Yagi, A . I. P o p o v a n d W . B . Person, / . Phys. Chem. 71 (1967) 2 4 3 9 ; J . Y a r w o o d a n d W . B . P e r s o n , / . Amer. Chem. Soc. 90 (1968) 594, 3 9 3 0 ; R . F . L a k e a n d H . W . T h o m p s o n , Spectrochim. Acta, 24A (1968) 1 3 2 1 ; S. G . W . G i n n , I. H a q u e a n d J. L . W o o d , ibid. p . 1 5 3 1 ; K . Y o k o b a y a s h i , F . W a t a r i a n d K . A i d a , ibid. p . 1 6 5 1 ; J. P . Kettle a n d A . H . Price, / . Chem. Soc.y Faraday Trans. II (1972) 1306.

CHEMICAL PROPERTIES OF THE HALOGENS

1199

TABLE 15. PROPERTIES OF SOME HALOGEN COMPLEXES IN SOLUTION

204

(a) Specific iodine complexes

Donor Benzene Naphthalene Methanol Ethanol Diethyl ether Diethyl sulphide Diethyl disulphide Ammonia Methylamine Dimethylamine Diethylamine Trimethylamine Pyridine

Formation constant, K (lmol-i)(20°C) 015 0-26 0-23 0-26 0-97 210 5-62 67 530 6800 6320 12,100 269

Heat of Charge-transfer band formation, λ e - A # ( k c a l m o l - i ) ιη & χ (m/x) Avi /2(cm~1) max 1-4 1-8 3-5 4-5 4-3 7-82 4-62 4-8 71 9-8 120 121 7-8

292 360 232 230 249 302 304 229 245 256 278 266 235

16,000 7250 13,700 12,700 5700 29,800 15,000 23,400 21,200 26,800 25,600 31,300 50,000

5100 4700 5700 6800 6900 5400 7200 4100 6400 6450 8100 8100 5200

(b) Halogen complexes generally 77-Donor, e.g. CeUe σ-Donor, e.g. amine

01-20 0-5-ca.l03

1-4 4-13

275-415 225-280

5000-15,000 -5000 3000-30,000 5000-8000

gas-phase molecules, the following decreases in frequency (Δν in cm - 1 ) are found for the stretching vibration of XY dissolved in benzene: Cl2, 31; Br2, 20; I2, 10; IC1, 28. These results support the view that the XY molecule is acting as a "sacrificial" σ-acceptor. For weak complexes, the proportional change in stretching force constant for the X-Y bond, Ak/k0, is given approximately by 2Δν/ν0 (where k0 and v0 refer to the unperturbed molecule), and on this basis the donor or acceptor capacities of different molecules have been compared168»197'201»204. However, for a relatively strongly bound complex, such asMe3N,XY, vibrational coupling has been shown to invalidate this simplified treatment of the force field214. The methods of vibrational spectroscopy have also been used to follow the course of ionization reactions such as 2py,IX v± [py2l]+ +IX2~

(X = Cl, Br or I; py = pyridine or y-picoline)

215

which occur in polar solvents . The enhanced intensity of certain modes with respect to infrared absorption has been the subject of numerous experimental and theoretical enquiries. For the halogen complexes it has been concluded216 that this effect, manifest in the halogen-halogen stretching vibra­ tion, is the outcome of electronic reorientation during the vibration, that is, of vibronic coupling between different electronic states; a similar effect has been described for hydrogenbonded complexes. (c) Since the total symmetry is likely to change on complex-formation, vibrational transitions which are forbidden in infrared absorption or Raman scattering in the free molecules may appear in the spectrum of the complex. Thus, the X-X stretching mode of 214 J. N . Gayles, / . Chem. Phys. 49 (1968) 1840. Acta, 23A (1967) 959, 2 5 2 3 ; S. G. W. Ginn and J. L. W o o d , 215 I. Haque and J. L. W o o d , Spectrochim. Trans. Faraday Soc. 6 2 (1966) 777. 216 H . B. Friedrich and W. B. Person, / . Chem. Phys. 44 (1966) 2161.

C.l.C. VOL II—PP

1200

CHLORINE, BROMINE, IODINE AND ASTATINE: A. J. DOWNS AND C. J. ADAMS

chlorine, bromine or iodine, normally inactive in infrared absorption, appears weakly in the infrared spectrum of the halogen in benzene solution. Although, in principle, the operation of the vibrational selection rules should afford a criterion of the geometry of a halogen complex, a simple interpretation is seldom free from ambiguity through the influence of vibronic coupling on band intensities and through other complications. Other physical properties used for studying complex-formation are as follows. Conductance1^ When the product of interaction of a halogen with a donor is appreciably ionic in character, complex-formation may be detected by conductance measurements. Thus, reactions such as that of IX and other halogens with pyridine (see above) give rise to solutions of appreciable conductivity, presumably containing the ions [py2X]+ and either halide or polyhalide anions. Solubility Studies™* A study of the solubility of iodine in normal solvents (giving a violet colour) was an important part of the work which led Hildebrand to his concept of "regular solutions"2*)?. All of these solutions conform reasonably well to the solubility equation for regular solutions RT\n(a2IN2)

=

V2i2$2-h)2

where a2 denotes the activity of the solid iodine referred to pure liquid iodine, V2 its liquid molal volume (extrapolated), φχ the volume fraction of the solvent and 82 a n d Sx the "solubility parameters" of iodine and the solvent. The experimental solubility of iodine in a complexing solvent is greater than the value predicted by this equation. Complexing solvents may also be distinguished by plots of log N2 against \jT which give a family of curves for regular solutions but curves with markedly different slopes for those solutions in which specific interaction occurs. Measurements of the solubility of iodine have also been used to determine formation constants for a number of complexes. For details of the solubilities of the halogens in aqueous and non-aqueous media, the reader is referred to Mellor's Comprehensive Treatise and Supplement32, Gmelins Hand­ buch33 or Linke's compilation of solubility data217. There have also been extensive studies of the adsorption of the halogens on materials such as charcoal, silica gel and magnesium oxide32; isotherms have been reported for several systems, and the influence of factors such as the preparation of the solid has also been investigated. Commonly there is but a fine distinction between purely physical processes of solution or adsorption on the one hand and overall chemical reaction on the other. Thus, under appropriate conditions iodine is reversibly adsorbed on certain metals or crystals of metal halides. Likewise, at low concen­ trations of iodine in aqueous solution the interaction with starch appears to be one of reversible adsorption; only at higher concentrations of iodine does the interaction give rise to a recognizable complex. The adsorption of bromine by graphite affords lamellar compounds by penetration between the carbon layers of the graphite; compounds with bromine contents up to that implied by the compositions C8Br or Ci6Br are thus obtained32»54. Dielectric Polarization Studies The product of interaction of a halogen and a donor may be more polar than either 217 W. F. Linke, Solubilities'. Inorganic and Metal-organic Compounds, 4th edn., Vol. 1, van Nostrand, Princeton (1958).

CHEMICAL PROPERTIES OF THE HALOGENS

1201

reactant. An estimate of the degree of polarization of the complex may then be obtained by measuring the dielectric constant of a solution of the complex in a non-polar medium. In general, the molar polarization of the halogen increases with the strength of the donor. Apparent dipole moments have also been calculated, notably for a number of iodine complexes168»218; values of4-12D have thus been reported for 1:1 complexes of iodine with various amines219. However, despite numerous attempts at correlation, no simple relation­ ship appears to exist between the scalar excess moments, i.e. ^(complex)-E^(components), and measures of acidic or basic character, e.g. formation constants219. Other Methods1™

Colorimetric measurements have confirmed that the heats of solution of iodine are generally greater in donor than in non-complexing solvents (e.g. cyclohexane < benzene < ethyl acetate < ethyl alcohol < pyridine). Investigations of the colligative properties and of the apparent molar volume of dissolved iodine have also been described. The magnetic susceptibility of benzene solutions of iodine is greater than expected from the normal additivity law, while very pronounced changes in magnetic susceptibility have been reported to accompany the formation of complexes of iodine or bromine with polycyclic aromatic bases such as perylene and pyrene, both in solution and in the solid state. Such adducts probably owe their unusual magnetic properties to the presence of significant concentrations of radical-ions, which would also account for their surprisingly high electrical conductivities. More recently the Ή nmr spectra of organic sulphide molecules or of methylpyridines have been used as indices to complex-formation with iodine or iodine monochloride220. Although the lifetimes of the complexes in solution are too short to permit the observation of more than averaged signals for a given proton, such measurements have been successfully applied to the determination of formation constants. The interaction of olefins and related hydrocarbons with molecular iodine supported on a celite or firebrick column has also been monitored by a gas-solid Chromatographie technique; the retention times of the hydro­ carbons afford measures of the relative formation constants for the olefin-iodine complexes, which follow a pattern similar to those for olefin-AgN03 complexes221. 3. Properties of solid complexes168»198»201 ~ 204

The theoretical interpretation of donor-acceptor interactions has stimulated numerous structural analyses based on X-ray diffraction studies of crystals of the halogen complexes. It has to be recognized, however, that the configurations found in the crystalline state are not necessarily the same as for individual complexes in solution or in the gas phase; only for especially stable complexes are discrete donor-acceptor units likely to be preserved in the different phases. In weak complexes the crystals typically consist of chains or sheets in which donor and acceptor units alternate in a regular way; a particular orientation of donor and acceptor may thus be favoured over others because it facilitates the attainment of a chain or layer structure, whereas the most stable orientation for a discrete complex in the gas phase may be quite different. Nevertheless, the following details of the crystalline structure are of especial interest: 218 S. N . Bhat and C. N . R. R a o , / . Amer. Chem. Soc. 9 0 (1968) 6008. 21 9 A . J. H a m ü t o n and L. E. Sutton, Chem, Comm. (1968) 4 6 0 . 220 j . Yarwood, Chem. Comm. (1967) 809; E. T. Strom, W. L. Orr, B. S. Snowden, jun., and D . E . Woessner, / . Phys. Chem. 71 (1967) 4017. 22i R. J. Cvetanovic, F. J. Duncan, W. E. Falconer and W. A . Sunder, / . Amer. Chem. Soc. 88 (1966) 1602.

s

O

N

Donor atom

400 400 400 400

2-37 2-37 2-37 2-37

1-65

3-20

O-Cl, 2-67

1,4-Dioxan, CI2

2-78 2-87 2-687 2-69

1-80

3-35

O-Br, 2-80

2CH3OH, Br2

S-I, S-I, S-I, S-I,

1-80

3-35

O-Br, 2-82

Acetone, Br2

Dibenzyl sulphide, I2 1,4-Dithian, 2I 2 1,4-Dithian, 2IBr 2Ph3PS, 3I 2

1-99 1-80

3-55 3-35

O-I, 2-81 O-Br, 2-71

1,4-Dioxan, 12 1,4-Dioxan, Br2

1-99

3-55

O-I, 2-57

1-84

3-45

N-Br, 2-84

2CH 3 CN, Br2

1,4-Dioxan, 2IC1

1-84

203 203 203 203 203 203

3-65 3-65 3-65 3-65 3-65 3-65 3-45

203 203

3-65 3-65

Van Covalent radius der Waals' sum (Ä) radius sum (Ä)

N-Br, 2-16

N-I, 2-26 N-I, 2-26 N-I, 2-57 N-I, 2-31 N-I, 2-92 N-I, 2 1 6 cation

N-I, 2-27 N-I, 2-30

Intermolecular contact (Ä)

Hexamethylenetetramine, 2Br2

Pyridine, IC1 Pyridine, IBr Pyridine, ICN 4-Picoline, 12 Phenazine, 12 Pyridine, 2I2

(a) σ-σ complexes Me 3 N, I 2 Me 3 N, IC1

Complex

I-I, 2-82 I-I, 2-79 I-Br, 2.646 I-I, 2-86

Cl-Cl, 2 0 2

Br-Br, 2-28

Br-Br, 2-28

Br-Br, 2-31

I-Cl, 2-33

Br-Br, 2-328

Br-Br, 2-43

I-Cl, 2-51 I-Br, 2-66 I-C — I-I, 2-83 I-I, 2-75

I-I, 2-83 I-Cl, 2-52

XY intra­ molecular bond length (Ä)

2-67 2-67 2-47 2-67

1-99

2-28

2-28

2-28

2-32

2-28

2-28

2-32 2-47 1-99 2-67 2-67

2-67 2-32

180 177-9 178-2 175

177

180

180

179-4

180

180

179 180

Bond Intra­ angle molecu­ lar bond i D - X - Y length in free XY 1 (°) (Ä)

Isolated Isolated Isolated Isolated

mol. mol. mol. mol.

Infinite chains

Infinite sheets

Infinite chains

Infinite chains

Isolated mol.

Isolated mol.

Isolated mol.

Isolated mol. Isolated mol. Isolated mol. Isolated mol. Infinite chains

Isolated mol. Isolated mol.

Structural type

TABLE 16. STRUCTURAL DATA FOR CRYSTALLINE HALOGEN COMPLEXES

d

S--I-Ilinear a I 2 molecule equatorial11 As in l,4-dioxan,2ICl e Crystal contains 2 Ph 3 PS- I 2 units linked by a "normal'' I 2 molecule (I-I = 2-73 Ä) f

Chains with equatorial dioxan-Br bonding8. Chains with O -Br-Br- O linear (see Fig. 15)a MeOH · Br 2 ·' MeOH units; linked into H-bonded structure11 As in l,4-dioxan,Br 2 a

Cl-I· -dioxan- I-Cl units; O I-Cllinear a

N tetrahedral; N · -Br-Br lineara N · ·Br-Br· -N linear0

Salt[py 2 I] + l3",2I 2 a

b

N- I-I lineara N--I-C1 ^ linear (within 3°)a N--I-Cl linear51 N--I-Brlinear a yC--N--I-CN linear» yC-N--I-Ilineara

Configuration

(b) σ-π complexes C 6 H 6 , Br 2 C 6 H 6 , Cl 2

Di(/?-chlorophenyl) telluride, 12

1 -Oxa-4-selenacyclohexane, IC1

1 -Oxa-4-selenacyclohexane, I2

1,4-Diselenacyclohexane, 2I 2 Selenacyclopentane, I2

C 6 H 6 -Br, 3-36 C 6 H 6 -C1, 3-28

Te-I, 2-947 2-922

Se-I, 2-630

Se-I, 2-755 3-708

Se-I, 2-83 Se-I, 2-76 3-64

3-65 3-50

4-35

4-15

415

415 415

2-70

2-50

2-50

2-50 2-50

Br-Br, 2-28 Cl-Cl, 1-99

1-1,3-85

I-Cl, 2-73

I-I, 2-956

I-I, 2-87 I-I, 2-91

2-28 1-99

2-67

2-32

2-67

2-67 2-67

175-8

174-8

180 179-4

Ί Infinite J chains

Isolated mol.

Isolated mol.

Isolated mol.

Isolated mol. Isolated mol.

Chains of alternate CÖHÖ and X2 molecules with X2 perpendicular to the plane of the CeHg ringa

Consists of G?-ClC6H4)2Tel2 molecules with axial I-Te-I bonds but with relatively short I· I intermolecular contacts1

Chair form with IC1 bonded to Se in the axial position 0

Structure resembles that of selenacyclopentane, I 2 e

Chair form with axial I 2 a Axial I2 molecule interacting more weakly with a second Se atom also axiala

a C. K. Prout and J. D. Wright, Angew. Chem., Internat. Edn. 7 (1968) 659; C. K. Prout and B. Kamenar, Molecular Complexes (ed. R. Foster), Paul Elek (Scientific Books) (1973). b T. Uchida, Bull. Chem. Soc. Japan, 40 (1967) 2244. c K.-M. Marstokk and K. O. Str0mme, Acta Cryst. Β24 (1968) 713. d O. Hassel, Acta Chem. Scand. 19 (1965) 2259. β C. Knobler, C. Baker, H. Hope and J. D . McCullough, Inorg. Chem. 10 (1971) 697. f W. W. Schweikert and E. A. Meyers, / . Phys. Chem. 72 (1968) 1561. e H. Maddox and J. D. McCullough, Inorg. Chem. 5 (1966) 522. h C. Knobler and J. D. McCullough, Inorg. Chem. 7 (1968) 365. 1 G. Y. Chao and J. D. McCullough, Acta Cryst. 15 (1962) 887; H. A. Bent, Chem. Rev. 68 (1968) 587.

Te

Se

1204

CHLORINE, BROMINE, IODINE AND ASTATINE: A. J. DOWNS AND C. J. ADAMS

(a) The general configuration of the donor-acceptor unit and the factors affecting the relative orientations of the donor and acceptor partners. (b) Interatomic distances relevant to the donor-acceptor contact and to internal bonds of the donor and acceptor molecules. Hence it may be possible to correlate the magnitude and importance of donor-acceptor forces in the crystal with the donor and acceptor strengths of the component molecules. Structural data so determined for σ-σ and σ-π complexes of the halogens are summarized in Table 16; the structures of some of the complexes are depicted in Fig. 15. From these results certain general features emerge. Thus, the crystalline σ-σ complexes are character­ ized by essentially linear D· · X-Y units, where D is the donor and X-Y the halogen acceptor; if X and Y are different halogen atoms then D is linked to the heavier atom. The D · · X contact is markedly shorter than the sum of the van der Waals' l-adii for the donor and acceptor atoms; it is invariably greater than the sum of the corresponding covalent radii, though for the strongest donors the margin becomes relatively narrow (as in l-oxa-4selenacyclohexaneJCl). With increasing donor strength the D · · X distance contracts in relation to the sum of the van der Waals' radii, while, in keeping with the anti-bonding character of the acceptor orbital, the intramolecular bond of the halogen molecule XY is attenuated. The order of donor strength Se > S > O and the order of acceptor strength 12 > Br2 > Cl2 may thus be deduced from the variations in interatomic distances given in Table 16. In the extreme case of pyridine with iodine the weakening of the I-I bond by charge transfer is so great that the bond is broken and the final product C5H5N,2I2 is best formulated as [(C 5 H 5 N) 2 I] + I 3 -,2I 2 ; here the cation is centrosymmetric and planar with two equal N-I distances of 2-16 Ä 222 . In the blue starch-iodine complex, which has been known for well over a century, some unusual structural details have come to light. The pertinent facts about this bluish-black complex are as follows: (i) it exhibits strong absorption in the 6250 Ä region; (ii) only amylose, the linear polymer fraction of the starch, forms the complex; (iii) the complexing unit and chromophore is not normally I 2 but Ι 2 · Ί ~ ; (iv) the iodine atoms are arranged in linear chains each of which occupies the central channel of a helix formed by the amylose unit; (v) there are about 3-9 glucose units per iodine atom223»224. The most recent measure­ ments of adsorption isotherms of the I 2 · · · I ~ system in the helical cavity of amylose have been interpreted225 on the assumption that the bound species can be expressed approxi­ mately as I 2 ,I~ b , where b varies between 0 and 1. A model for the structure is believed to have been found in the complex HI 3 ,2C 6 H 5 CONH 2 , wherein the benzamide molecules are associated via hydrogen-bonding and stacked in such a way as to leave long channels in which nearly linear I3 _ ions are aligned226. There is evidence of strong attraction between successive I 3 _ ions with d(l · · -I) = 3-80 A as against intramolecular distances of 2-90 and 2.96 Ä. There are two such polyiodide chains in each channel probably cross-linked by hydrogen bonds formed by the HI3 protons; the internal diameter of the amylose helix does not appear to be large enough to accommodate more than one chain. Interaction between the 13 ~ ions is believed to be responsible, not only for the blue colour, but also for the esr

222 o . Hassel and H. Hope, Acta Chem. Scand. 15 (1961) 407. 223 R . Bersohn and I. Isenberg, / . Chem. Phys. 35 (1961) 1640. 224 B . S. Ehrlich and M. Kaplan, / . Chem. Phys. 51 (1969) 603. 225 c . L. Cronan and F. W. Schneider, / . Phys. Chem. 73 (1969) 3990. 226 j . M . Reddy, K. Knox and M. B. Robin, / . Chem. Phys. 40 (1964) 1082.

xfi

(·)

*x

a°'

0 ? r Oo oc

CH, CH3

(b)

•Οζ-

CH, CHL3 •x / C

A x

Br

\

/

Br

Br

Br O

>-o-^ V - 0 ^

Ο'ΒΓ

ΟΌ

f-o-o

oc

c sin j3

Oßr OC FIG. 15. Schematic diagrams of the structures of crystalline halogen complexes: (a) 1,4-dioxan, Br2i (b) acetone, Br2; (c) 2CH3OH, Br2; (d) CÖH 6 , Br2 projected along a axis. [Reproduced from Molecular Complexes (ed. R. Foster) by kind permission of Paul Elek (Scientific Books) Ltd. and C. K. Prout and B. Kamenar.]

1205

1206

CHLORINE, BROMINE, IODINE AND ASTATINE: A. J. DOWNS AND C. J. ADAMS

spectrum given by the starch-iodine complex, and on the basis of such evidence the intriguing proposal has been made 223 that the rows of iodine atoms function as one-dimensional metals. However, Mössbauer resonance experiments focused on the 129I nucleus give unambiguous notice that the I 3 ~ ions, whether in solid Csl3, HI 3 ,2C 6 H 5 CONH 2 or amylose-I3 ~, retain non-equivalent iodine atoms224. Accordingly the notion of a one-dimensional metal must be treated with some reserve. Other polyhydric compounds such as amylopectin, formed by the hydrolysis of starch (e.g. by enzyme action), also give highly coloured complexes with iodine. Differences in the spectroscopic properties of such complexes have been exploited for analytical measurements. Charge transfer may ultimately proceed to the reduction of the halogen and simultaneous oxidation of the donor. Thus the interaction of a compound of the type R2Se or R2Te (R = organic group) with a halogen may produce, not an adduct, but the corresponding dihalide, a molecular compound with a nearly linear X-Se-X or X-Te-X skeleton. However, the crystal structure of one such compound (/?-ClC6H4)2TeI2 still exhibits certain idiosyncracies suggestive of donor-acceptor interaction: one Te-I bond is significantly longer than the other, while the intermolecular I · · · I packing distance in the line of the Te-I bonds is 3-85 Ä, as against 4-30 Ä for the sum of the van der Waals' radii202. It appears, then, that by variation of the donor it is possible to move in gradual steps from a well-defined molecular adduct such as dibenzyl sulphide,I2, where the I-I interaction is relatively strong, through an intermediate stage, represented for example by selenacyclopentane,I2, where the I-I interaction is relatively weak and some I-Se· · -I bonding exists, to a compound such as (/?-OC6H4)2TeI2, where the I · · · I interaction is very weak and I-Te-I bonding is relatively strong. In those complexes where the donor-halogen interaction is relatively strong, discrete molecular units are usually identifiable. By contrast, crystals of the weaker complexes frequently contain chains in which donor sites are linked by linear D · · X-X · · · D bridges. In many respects these bridges are comparable with hydrogen bonds, an analogy well displayed in the crystalline complex Br2,2CH3OH (Fig. 15): each oxygen atom is here coordinated to three neighbouring oxygens by two hydrogen bonds and one O · · · Br-Br · · · O bridge. Bridging halogen molecules are likewise prominent in the crystalline σ-π complexes X 2 ,C 6 H 6 (X = Cl or Br) on the premises of two-dimensional crystallographic analysis assuming the centrosymmetric space group C2/l3. The structure consists of infinite chains of alternate benzene and halogen molecules with the latter placed perpendicular to the planes of the benzene rings along a common sixfold axis and equidistant from successive pairs of benzene molecules. The benzene-halogen distances are slightly shorter than the sum of the corresponding van der Waals' radii, but the X-X distances are not measurably different from those in the free halogen molecules. In the face of these conclusions, the infrared spectra of single crystals of Br2,C6H6 are consistent, not with the centrosymmetric disposition of benzene and bromine molecules required by the space group C2ft3, but rather with a structure having alternate long and short benzene-bromine distances227. This apparent contradiction is probably a result of disorder in the crystals. In all probability there are but marginal differences of energy between the different configurations of such a weakly bonded complex; in solution the complex may therefore assume one or more highly labile configurations quite unrelated to the C 6 H 6 - · ·Χ 2 axial unit characteristic of the crystal. 227 w . B. Person, C. F. Cook and H. B. Friedrich, / . Chem. Phys. 46 (1967) 2521.

CHEMICAL PROPERTIES OF THE HALOGENS

1207

Measurements of halogen nqr frequencies112-228 and of the Mössbauer effect in the ^I nucleus229 have also been applied to solid complexes of the halogens. Approximate charge distributions calculated from the nqr frequencies of complexes of amines with Br2, I2, IBr or IC1 clearly signify the transfer of charge from the amine to the halogen molecule in the electronic ground state of the complex. The transferred charge appears to go mainly to the halogen atom Y in the unit D · · -X-Y, with the result that in complexes of the homonuclear halogens the atom X acquires a partial positive charge. Such a charge distribution correlates well with a scheme of delocalized σ-bonding encompassing the three centres D · · -X-Y (compare Fig. 4 and p. 1560), and according to which the trihalide ions, amine-halogen complexes and [(amine)2halogen]+ cations constitute an isoelectronic series. On the other hand, the nqr properties of the solid C6H6,Br2 complex vouch for the impli­ cation of the X-ray data that the degree of charge-transfer is here very small230. 12

4. Factors influencing the stabilities of halogen complexes168 » 196 » 197 » 201 ' 204

Formation constants and related thermodynamic parameters have been measured for a large number of halogen complexes in solution. Combined with the interatomic dimensions of those solid complexes which have been subjected to structural analysis, such results serve as a basis for evaluating donor-acceptor interactions as influenced, for example, by changes in the acceptor, the donor or the environment. The consensus of the evidence is that the relative acceptor strengths of the halogens with respect to a given donor follow the sequence IC1 > BrCl > IBr > I 2 > Br2 > Cl2. This order has been found to apply to interactions with donors as diverse as aromatic hydrocarbons and halide ions, though it varies somewhat with the choice of donor; certainly the response of the formation constant to changes in the donor varies in magnitude from acceptor to acceptor, being greatest for the strongest acceptor. It is difficult to compare the acceptor strengths of the halogens with those of other Lewis acids because of the wide variation in conditions which have been used in studying different kinds of complex. However, aromatic hydrocarbons form adducts with sulphur dioxide intermediate in stability between the complexes formed with bromine and with chlorine, while iodine and 1,3,5-trinitrobenzene appear to be comparable in strength, if not in electronic action, as acceptors. Analogous considerations suggest that, with respect to a given halogen, donor capacity increases in the sequence benzene < alkenes < polyalkylbenzenes « alkyl iodides « alcohols « ethers « ketones < organic sulphides < organic selenides < amines. Again some latitude must be allowed for such influences as the acceptor strength of the halogen or steric factors, which may assume importance in many situations. For example, complexes formed by l,3,5-tri-ter/-butylbenzene are much less stable than those formed by mesitylene presumably because of the bulk of the alkyl substituents, which inhibits the close approach of the donor and acceptor moieties. The halogens thus emerge as "soft acids". Indeed, iodine is favoured as a reference soft acid for evaluating the relative coordinating power of different donor species, including a wide range of non-aqueous solvents231. In this connection 228 G. A . Bowmaker and S. Hacobian, Austral. J. Chem. 2 2 (1969) 2047. 229 c . I. Wynter, J. Hill, W. Bledsoe, G. K. Shenoy and S. L. Ruby, / . Chem. Phys. 50 (1969) 3872. 230 H . O. Hooper, / . Chem. Phys. 41 (1964) 599; D . F. R. Gilson and C. T. O'Konski, ibid. 48 (1968) 2767. 231 R. S. D r a g o and K. F. Purcell, Non-aqueous Solvent Systems (ed. T. C. Waddington), p. 225, Aca­ demic Press (1965).

1208

CHLORINE, BROMINE, IODINE AND ASTATINE.* A. J. DOWNS AND C. J. ADAMS

use has been made of the Drago-Wayland relationship232, -Mf=

EAEB+CACB

which resolves the enthalpy of adduct-formation, Δ//, into parameters E and C representing the susceptibilities of the acid (A) and base (B) to engage in electrostatic interaction (£) and covalent bonding (C). The relationship expresses the generality that all donor-acceptor interactions are composed of some electrostatic and some covalent properties; to this extent so-called "hardness" and "softness" are not mutually exclusive characteristics. The selection of iodine as a reference acid entails the assignment of equal values (unity) of EA and C A 233 . The thermodynamic properties designating the formation of halogen complexes show significant variations as the solvent is changed. The enthalpy of complex-formation measured in solution A//°(soln) is related to that measured in the gas phase A//°(g) by A/T(soln) =

Δ^°(8)-Δ^301ν(Χ2)-Δ^301ν(ϋ)+Δ^801ν(0,Χ2)

where Δ// δ0ΐν denotes the enthalpy of solvation; analogous relationships exist for the free energy and entropy changes attending complex-formation. For a system such as MeCONMe2 + l2 ^ MeCONMe2,l2> only in dilute solution in an inert solvent like carbon tetrachloride or a saturated hydrocarbon is the net solvation energy small compared with the enthalpy of adduct-formation, and only for such solutions is it safe to assume that the enthalpy measured closely corresponds to A/f°(g). Even with inert solvents the thermo­ dynamic properties of complex-formation vary somewhat from solvent to solvent, a circumstance which can commonly be attributed to minor differences in activity coefficients of the reactants and complex in the different media; variations in the degree of aggregation of one or more components represent another potential source of solvent-dependent thermodynamic properties234. A much larger variation results, however, with polar solvents which themselves have some capacity to act either as donors or acceptors. Under these conditions the solvation energies and the balance of such energies may approach, or become relatively large compared with, the measured enthalpies of adduct-formation; this is found to be the case even with a medium as weakly solvating as dichloromethane232. Solutions in inert solvents are amenable to "regular" solution theory: for example, one such treatment relates the formation constant K of a. complex to the solubility parameter 8S of the solvent by the equation logK=

a+b8s

where a and b depend only on the properties of the donor and acceptor235. More generally, however, the effect of solvent perturbation is incompletely understood; it is not even convincingly established that the interactions stabilizing the complex itself are the same in the gas phase as in solution2**4. It is a general rule that the intensity of the charge-transfer band of a complex in the vapour phase is considerably lower than that of the same complex in solution. The reason for this phenomenon is again not clear. One suggestion is that the solvent cage around the complex in solution confines it to the extent that it is under some pressure, resulting in enhanced 232 R . s . D r a g o , Chem. in Britain, 3 (1967) 516. 233 R . s . D r a g o and B . B . Wayland, / . Amer. Chem. Soc. 87 (1965) 3571; O. W . Rolling, Inorg. 8 (1969) 1537. 234 w . Partenheimer, T . D . Epley and R. S. D r a g o , / . Amer. Chem. Soc. 9 0 (1968) 3886. 235 p . v . Huong, N . Platzer and M . L . Josien, / . Amer. Chem. Soc. 91 (1969) 3669.

Chem,

CHEMICAL PROPERTIES OF THE HALOGENS

1209

overlap of donor and acceptor orbitals and hence in an increased transition moment for the charge-transfer absorption. This is consistent with the observed effects of externally applied pressure on the spectra of π-π complexes in solution204. 5. Nature of the interactions in halogen complexes 168» 196 ~ 204

As a rule donor-acceptor interactions are intermediate in character between ordinary van der Waals' contacts and normal covalent bonds. Such interactions, which include hydrogen bonds, may be viewed as the first steps of bimolecular nucleophilic displacement reactions. To this extent there is a close relation between the formation of molecular adducts and certain types of reaction of the elementary halogens. Numerous theories have been developed concerning the nature of donor-acceptor interactions; the views of the early theorists have been surveyed elsewhere in some detail 236 ' 237 . According to the various approaches, the interactions have been described chemically by such phrases as "saturation of residual affinities", "exaltation of valency", "secondary acid-base interactions" or "facecentred bonding"; physically in terms of "adhesion by stray feeler lines of force", "adhesion by attraction of positive and negative patches in molecules", "charge-sharing" or "chargetransfer"; and quantum mechanically by reference to "complex resonance", "no-bonddative-bond resonance", "filling of anti-bonding orbitals" or "interaction of the highest occupied orbitals (of one component) with the lowest vacant orbitals (of the other)" 202 . These phrases are all useful and complementary, emphasizing different aspects of intermolecular interactions. Some phrases emphasize the directional character of strong intermolecular interactions; or the intermediate length and strength of intermolecular as compared with intramolecular bonds; or the similarity of intramolecular and intermolecular bonds; or the creation of formal charges and "expanded octets"; or the fact that, like intramolecular bonding, intermolecular bonding is essentially an electrostatic phenomenon. An adequate interpretation of donor-acceptor interactions, applicable to the halogen complexes, must account satisfactorily for the appearance of the intense absorption bands which are characteristic of these adducts, and for the variations in the frequencies and intensities of the bands with changes in the nature of the components. In addition it should be consistent, insofar as it is applicable to the prediction of the orientation of the compon­ ents with respect to each other, with the thermodynamic properties, vibrational spectra and dipole moments of complexes in solution and with the crystal structures of solid adducts. In this sense the theoretical basis of the interaction of donor and acceptor molecules first given in 1952 by Mulliken, the so-called charge-transfer theory of complex-formation238, has gained the widest acceptance. According to Mulliken's model the donor-acceptor complex D,A in its ground state is best represented as a resonance hybrid or combination of a "no-bond" wavefunction ^0(D,A) and one or more "dative-bond" functions such as ^ 1 (D+-A _ ). The no-bond function includes the electronic energy of the component molecules plus terms representing the effect of dipole interactions, disperson forces, hydrogen-bonding and other inter­ molecular forces. The dative-bond functions represent states where an electron has been transferred from the donor molecule to the acceptor, introducing electrostatic interactions and forming a weak covalent link between the resulting radical ions. In quantum-mechanical 236 L. J. Andrews, Chem. Rev. 54 (1954) 713. 237 L. E. Orgel, Quart. Rev. Chem. Soc. 8 (1954) 422. 238 R. s. Mulliken, / . Amer. Chem. Soc. 72 (1950) 600; ibid. 74 (1952) 811; / . Phys. Chem. 56 (1952) 801; / . Chem. Phys. 23 (1955) 397; Rec. trav. chim. 75 (1956) 845.

1210

CHLORINE, BROMINE, IODINE AND ASTATINE! A. J. DOWNS AND C. J. ADAMS

terms the wavefunction for the ground state is approximated by φΝ = αψ0φΛ) + οφι(Ό+-Α-)

(1)

For a weakly bonded complex a > b, a and b the mixing coefficients being related by the normalization condition a2 + b2 = 1. Excited states, with a dative structure as the main contributor, have the same form with the coefficients varied to give predominance to the dative-bond contribution. Thus the wavefunction of such a charge-transfer state is given by ψΒ = a*
(2)

wherein a* > ft*. The coefficients a* and Z>* are determined by the quantum-theory require­ ment that the excited-state wavefunction be orthogonal to the ground-state function: $ΦΝφΕατ = 0. It follows that a* & a and Z>* ^ b. The results of the resonance interaction are then: (i) Admixture of the charge-transfer state gives a ground state for the complex which is more stable than that represented by any of the component wavefunctions. (ii) Absorption of light induces a transition from the ground to the excited state with transfer of an electron from an orbital largely associated with the donor to an orbital largely associated with the acceptor. This accounts for the charge-transfer band character­ istic of the complex. Intensity of the Charge-Transfer Band The transition moment for the absorption μΕΝ is given by μεΝ = —βίφεΣηφΝατ

(3)

Γ| being the position vector of the /th electron. Hence it may be deduced that μΕΝ

= α*δ{μ\ - μο) + (aa* - 66*)Οοι - Spo)

(4)

where μι = - e j ^ l r ^ r f r , /*0 = -«#Ο Σ Γ <Λ>*·>/*οι = ~βΙΦ\^ιΦΦ and S = ίφ0φ\ατ. There are therefore two contributions to the transition moment. The first term in μχ — μο is proportional to the dipole moment of the transferred electron and the hole it leaves behind and is related to the stabilization of the ground state through the coefficient a*b. If this term is to contribute to the intensity of the charge-transfer band, a*b must be greater than zero, implying some region of overlap between the orbitals of D and those of A; the greater the overlap the more intense the band. However, Mulliken has pointed out that, even if a*b is very small or zero, so that there is negligible stabilization of the ground state, a chargetransfer band may still be observed through the influence of the second term in equation (4). This provides a plausible explanation of the charge-transfer absorption observed in systems such as iodine in n-heptane or cyclohexane, wherein complex-formation in any sense other than statistical collision-pairing is improbable. In a chance collision encounter between the hydrocarbon and halogen molecules the ground-state potential energy curve is unlikely to exhibit a minimum—that is, no recognizable complex is formed—but the orbitals of the components may overlap adequately to give a mixing of the non-bonded and the chargetransfer states, with a substantial transition moment. The expression "contact charge-trans­ fer" has been coined to describe the corresponding absorption which appears, for example, in iodine-heptane mixtures near 260 τημ. Even in the cases where a definite complex can be identified, collision charge-transfer, as well as complex charge-transfer, contributes to the total intensity of the observed absorption. In a series of related complexes, the portion of the

CHEMICAL PROPERTIES OF THE HALOGENS

1211

intensity arising from the complexed donor-acceptor pairs progressively increases and that contributed by the collision-pairs decreases as the complexes become more stable. Calcula­ tions suggest that as much as three-quarters of the intensity of the observed charge-transfer band of the iodine-benzene system may be derived from collision-pairs200. In all probability there is a continuous gradation between the two extremes represented by a true complex and a contact-pair, corresponding to a wide variation in the "stickiness" of collisions responsible for charge-transfer absorption. There may also be some flexibility in respect of those orientations of the components which are required in order that absorption might occur. A mixing of the charge-transfer transition with internal transitions of the component molecules, particularly those of the donor, can also contribute additional intensity to the charge-transfer absorption. According to Murrell199, such mixing even accounts for the so-called "contact charge-transfer" phenomena. Energy Terms Application of the variation principle gives for the ground-state energy E associated with the total wavefunction of the complex (E0-E)(E1-E) = (H0i-ES)2

(5)

Here E0 = $ψοΗψ0ατ is the energy associated with the structure D,A, E\ = $φιΗψχατ is the energy associated with the charge-transfer state D + — A ~, /70i = ΙΦ^Ηφιάτ is the interaction energy due to the mixing of ^o a n ^ Φι, Η is the total exact Hamiltonian for the entire system, while, as before, S = ΙΦαΦχάτ denotes the overlap of the functions ^o and φ\. From equation (5) it follows that E=

Eo-iHoi-ESMEi-E)

There are two solutions of this equation for E. One corresponds to the final ground state (EN) and the other to the final excited state (EE). Because the energy of interaction between a halogen and donor molecule is typically small and E\ — E is relatively large, EN approxi­ mates to E0. Then EN = EQ- (Hoi - EoSy/iEi - E0)

(6a)

while the corresponding energy of the excited state is EE = £i+(tfoi - EiSWEi - £b)

(6b)

Further, the mixing coefficients a and b of the ground state and a* and b* of the excited state are given by - = -(ffoi-EoMfa-Eo) a

% = -(#oi-£iS)/(£i -£o) cr

(7)

Equation (6a) emphasizes that the energy of the ground state has contributions both from classical intermolecular forces, electrostatic in origin, through E0 and from covalent interactions principally embodied in the second term. The partitioning of the energy into electrostatic and covalent parts provides a theoretical justification for the Drago-Wayland relationship (see p. 1208). The energetics of complex-formation are illustrated schematically in Fig. 16. The stabilization energy of the ground state arising from the mixing of ψ0 and ψΪ9 Ε0-ΕΝ, is comparable with, but generally larger than, the heat of formation of the complex ΔΗ; for the benzene-iodine complex, for example, E0 — EN «* 1-3 kcal mol - 1 . As a rule S and Η0χ are also relatively small; a and a* are close to unity and thus large compared with b and b*9

1212

CHLORINE, BROMINE, IODINE AND ASTATINE: A. J. DOWNS AND C. J. ADAMS

(a)

D + e+ A

(b)

E

Energy

i

/

\

Intermolecular separation

EN

»►

FIG. 16. The relations between the energies of the ground (N), non-bonded (D,A), chargetransfer (D + — A~), and excited (E) states and the distance between the donor (D) and acceptor (A) molecules in a charge-transfer complex. E8 is the energy of the infinitely separated donor and acceptor; Es — E0 = stabilization energy due to classical intermolecular interactions; and E0—EN = stabilization energy due to electron delocalization.

while Εχ - E0 is a relatively major energy term, being about 180 kcal mol ~i for the benzeneiodine complex, for example. If ΙΌ is the ionization potential of the donor and EA the electron affinity of the acceptor, Fig. 16 indicates that these quantities are related to the frequency v of the charge-transfer transition by hv = / D _ £ A _ A

(8)

Δ being the difference between the binding energies of the components in the ground and the excited states. The most important components of Δ are the coulombic energy Ec of the charge-transfer state D + - A - and the resonance energies E0 — EN and EE — Ei arising from the mixing of the charge-transfer and non-bonded states. Good linear correlations are found between the charge-transfer frequencies of different iodine complexes and the corresponding ionization potentials of the donor molecules. Similarly, for a particular donor the frequency of the charge-transfer absorption has been shown to be proportional to the electron affinity of the acceptor. However, the slopes of the straight lines relating hv to ΙΌ for the iodine complexes are invariably less than unity, implying that Δ varies with the nature of the donor. Account of this is taken in the following approximate relation which can be justified theoretically for a set of closely related weak complexes of a single acceptor: hvX / D - C i +

C2 /D-CI

(9)

Here C\ & EA—EC + (ES — E0)9 the term (Es — E0) representing the interaction energy of the donor and acceptor in the formation of the no-bond structure defined by φο, while C
CHEMICAL PROPERTIES OF THE HALOGENS

1213

either C\9 with its dependence on the coulombic energy Ec, or the resonance integrals which determine C2 to remain constant for all donors, even with the same acceptor. The shift to higher energies in the visible absorption band of iodine as a result of complexformation corresponds to about 1 ·5Δ#. When the iodine molecule, paired off with a donor partner, is excited by the absorption of visible light (au <- ng), its suddenly swollen size corresponding to the occupancy of the strongly anti-bonding au orbital introduces an exchange repulsion between it and the donor. It is suggested that the repulsion energy, which should be greater the more intimate the donor-acceptor contact, is added to the usual energy of the excited iodine molecule, thereby increasing the frequency of absorption204. Correlation of the visible absorption spectrum of iodine in a wide range of solvents with the ionization potentials of the solvents has recently provided a distinctive classification of the solvents, the variations being attributed to the influences of charge-transfer and contactcharge-transfer interactions239. Curiously, however, a recent examination of the spectrum of the strong complex Et2S,I2 in the vapour state shows no blue shift in the visible band due to iodine240. No good explanation of this behaviour has yet been given. The energies of the two states represented by φ0 and φχ depend upon the relative orienta­ tions of the donor and acceptor; only if ψχ is much more sensitive to such effects than φ0 will the charge-transfer interaction tend to dictate the structure adopted by the complex. Because the energy of the charge-transfer state D + —A~ may in certain circumstances be a sensitive function of the nature of the solvent, there is not necessarily any relation between the structure of the complex in the solid and solution phases. According to Mulliken's "overlap and orientation" principle, maximum charge-transfer interaction is achieved when the relative orientations of the donor and acceptor molecules provide maximum overlap of thefilleddonor and vacant acceptor orbitals. This principle is neither proved nor disproved by determinations of individual crystal structures: it is a theoretical principle applying to idealized systems involving no intermolecular forces other than those due to charge-transfer. In reality the configuration of a complex is also affected, and in some cases largely controlled, by electrostatic attraction and exchange-repulsion terms associated with the no-bond structure. Attempts by Mulliken to predict the structures of certain halogen complexes, e.g. C5H5NJ2 and C6H<5,I2, on the strength of the overlap and orientation principle, led to conclusions at variance with the structures subsequently identified in the solids; presumably this discord simply reflects the conflicting and subtle requirements of the factors which minimize the total energy of such a system. More Generalized Treatments Although the principles of Mulliken's model of charge-transfer interactions have come to be generally accepted, in its simplest form the model is often found to be imperfect. A more general form for the wavefunction of the ground state of a donor-acceptor complex is φΝ = a « A o ( D , A ) + S W i i ( D + - A - ) + E c i 0 2 i ( D - - A + ) + · ■ ·

(10)

wherein more than one charge-transfer state is incorporated. Subject to the overall geometry of the complex, the coefficients b% and ci differ from zero only when the corresponding wavefunctions ψα and 02y belong to the same group-theoretical species as ψ0. A full treat­ ment should also take account of excited states of the donor and acceptor molecules by inclusion of no-bond functions such as ψ0(Ό*,Α) and of charge-transfer functions such as 239 E. M. Voigt, / . Phys. Chem. 72 (1968) 3300. 240 See for example M. Kroll, / . Amer. Chem. Soc. 90 (1968) 1097.

1214

CHLORINE, BROMINE, IODINE AND ASTATINE: A. J. DOWNS AND C. J. ADAMS

φι(Ό + — A -*). In many cases only the first term and one or two terms of the first summation are important, but the possibility of a significant contribution from a charge-transfer state D — A + in which the halogen assumes the donor function cannot be excluded. There has been considerable debate about the precise nature and relative importance of classical electrostatic forces (e.g. dipole-dipole and dipole-induced-dipole) as against charge-transfer interactions. It has recently been shown, for example, that quadrupoleinduced-dipole forces may be comparable in magnitude with charge-transfer interactions in benzene-iodine and related complexes241. While some authors have argued the importance, if not preponderance, of coulomb and polarization forces in the formation of weak complexes, others have sought to invoke London dispersion forces as making major contributions to the energies of formation of such complexes. However, although dispersion forces neces­ sarily contribute to the stability of complexes in the vapour phase, their effects are approxi­ mately cancelled out in the formation of complexes in solution. A recent analysis concludes that electrostatic, charge-transfer and exchange-repulsion interactions are all important in describing molecular complexes241; in the weakest complexes the extent of charge-transfer action may have been overestimated by Mulliken's conventional description, electrostatic forces being pre-eminent. Nevertheless, such considerations do not conflict with the principles of the Mulliken treatment, implying rather the inadequacy of its quantitative use in a highly simplified form. Alternative descriptions have been based on delocalized molecular orbitals extending over the complex as a whole. A complete account would require all-electron, self-consistentfield MO calculations which have not yet proved tractable for the halogen complexes. However, simplified LCAO-MO approximations have been shown to be compatible with the resonance-structure approach developed by Mulliken, and to underline the similarity between a polyhalide ion like I3 ~ and a strong molecular complex like R3NJ2 (see p. 1560). Thus, one such approach assigns analogous wavefunctions to the ground states of R3NJ2 and of I3 -, e.g. MR3NJ2) « a«Ao(R3N,Ia-Ib) + ^ ( R 3 N + - I a , I b - ) + # ( R 3 N + I a - I ö )

(11)

Here the coefficients c and d are unequal; according to nqr measurements112»228 c > d. The distinction between I3 ~ and R3NJ2 hinges on the relative magnitudes of a, c and d: if d is assumed small, a = c for the symmetric complex I 3 -. A similar situation in R3NJ2 would correspond roughly to a structure R3N+* · · · I a · · · Ib -* with "half-bonds" between N and la and between I a and I b . However, the N-I and I-I distances in the solid complex Me3N,I2 suggest that c is in fact significantly smaller than a. 6. Intermediacy of halogen complexes in chemical reactions202*204 In some complexes two isomeric forms are possible differing strongly in the extent of charge-transfer and often, at the same time, in molecular configuration. Usually only one of the two forms is stable in any one environment, but a change of environment (e.g. from a non-polar to a polar solvent) may favour a change from one to the other form. Such iso­ meric forms have been designated "outer complexes" (usually loosely bound and with little charge-transfer) and "inner complexes" (usually with more or less complete charge-trans­ fer). The transformation of an outer halogen complex to an inner complex involves dis­ sociation of the halogen molecule, as illustrated in Fig. 17 and by the following examples: 241 M. W. Hanna, / . Amer. Chem. Soc. 90 (1968) 285; M. W. Hanna and D. E. Williams, ibid. p. 5358; J. L. Lippert, M. W. Hanna and P. J. Trotter, ibid. 91 (1969) 4035.

1215

CHEMICAL PROPERTIES OF THE HALOGENS

+

(X) -

•O-O"

φ-οσ -O" - §>0

+

-

O'

Inner complex y Non-polar medium Energy

Polar medium

Reaction coordinate FIG. 17. The formation of "inner" and "outer" complexes of a halogen with a trialkylamine molecule: bimolecular nucleophilic substitution at one of the halogen atoms. Outer complex (i)

C5H5N + I 2 ^=±C 5 H 5 N, I2

(ii)

NH 3 -M 2 ^=±H 3 N,l 2

C5H5N ~Slow aq. NH 3

Inner complex [(C5H5N)2I]+I[H3NI]+I-

1[NH 3

(iii)

TMH + I 2 Ξ = ϋ ΤΜΗ,Ι2 ||TMH 2TMH,I2

Polar solvent

NH2I + NH4++ Γ

H

(TMH=Me2NNMe2)242

Similar reactions with the formation of radical-ions probably occur in the interaction of iodine with donors like perylene or carotene. (iv)

MR34- X2 Ξ = ± : R3M, X2

[R3MXl+X ^±R 3 MX 2 [R3MX]+X3'

(M=P, As or Sb; R = CH3or C6H5)243

242 D . Romans, W. H. Bruning and C. J. Michejda, / . Amer. Chem. Soc. 91 (1969) 3859. 243 See for example S. N. Bhat and C. N . R. Rao, / . Amer. Chem. Soc. 88 (1966) 3216.

1216

CHLORINE, BROMINE, IODINE AND ASTATINE: A. J. DOWNS AND C. J. ADAMS Inner complex

Outer complex aq.

(v)

[H 2 OX]+X~

IK (vi)

H,S + I 2

(vii)

R2Y + X 2 :

HOX + H3O++ X [H 2 SI]+I- (ref. 244) CR 2 YX;HX-V± R 2 Y X 2

R 2 Y,X 2

(Y = Se or Te; R = organic group) 202,245

Polar solvent

(viii) (ix)

+ Br,

VIh-Br, C

/ x

Slow Br,

\f==\+lix-

L \^^y ^x +

\c/

l>Br

>\ J

Br-

The stages involve a bimolecular nucleophilic displacement reaction, comparable with a Waiden inversion, at one of the halogen atoms; the formation of a symmetrical complex like I3 ~ corresponds to a transition state, sometimes referred to as a "middle complex"204, intermediate between the outer and inner complexes. The facility with which such dissociative reactions occur depends upon the natures of the halogen molecule, the donor atom and the substituents borne by the donor atom; the forma­ tion of an inner complex is generally a mark of strong donor-acceptor interaction. In contrast with acceptors like the hydrogen halides, which are highly prone to undergo dissociation, the homonuclear halogen molecules are relatively much more reluctant to function in this way. In other words the halogen molecules commonly function simply as sacrificial acceptors, whereas with the hydrogen halide molecules the sacrifice goes virtually to completion; alternatively we may say that the halogen atom X in the molecule X 2 is a relatively poor "leaving group" with respect to nucleophilic displacement at the second atom. However, the relative stabilities of inner and outer complexes are highly susceptible to the influence of environment. In particular, the ionic inner complex is stabilized relative to the outer complex (i) by a polar solvent both through the action of classical electrostatic forces (dependent on properties such as the dielectric constant of the medium) and through the action of the solvent as an auxiliary donor or acceptor, and (ii) by ionic crystal-formation resulting in an increase in coulomb energy per ion-pair as compared with the isolated inner complex. The influence of the solvent is shown qualitatively in Fig. 17: in a non-polar medium the inner complex, while it may have a potential minimum, is typically too high in energy to be formed, while the outer complex is more or less stable as a loose complex; on the other hand, in a sufficiently polar medium the relative stabilities of the two states may well be reversed. Such environmental assistance of the heterolytic dissociation of halogen molecules probably plays a central role in the reactions of the halogens in the condensed phases, and particularly in polar solvents such as water. Some of the more important examples of such reactions are considered briefly in the following sections. 244 j . Jander and G. Turk, Ber. 95 (1962) 881. 245 G. C. Hayward and P. J. Hendra, / . Chem. Soc. (A) (I960) 1760.

CHEMICAL PROPERTIES OF THE HALOGENS

1217

Reaction of the Halogens with Water and Hydroxide Ions (see pp. 1188-95 and Table 13) The rate of hydrolysis of chlorine, bromine and iodine follows the order Cl2 < I2 < Br2. In all three cases the intermediate is thought to be the outer complex [X2,OH] - or X 2 ,OH 2 (depending on the conditions), so that the overall reaction rate depends upon the formation of this species and its subsequent transformation to an inner complex: X2 + OH- ^[X 2 ,OH]- - ^ X + X O H

The stability of the transition state presumably decreases in the order I 2 > Br2 > Cl2, but the subsequent decomposition depends upon the X-X bond strength, which increases in the order I 2 < Br2 < Cl2. Understandably this renders unpredictable the overall kinetic pattern. The intermediacy of the species [I2,OH] ~ has been corroborated by studies of alkaline solutions of iodine, which suggest that k\ is about 100 times greater than &2 24(5. Chlorine hydrate is of historical interest as the solid phase originally thought to be solid chlorine but shown (in 1811) by Davy to contain water247; it was later assigned the formula Cl2,10H2O by Faraday 248 . Subsequent reports of the crystalline hydrates formed at temperatures close to 0°C gave compositions ranging from C12,6H20 to C12,8H20, but, according to a more recent investigation249, passage of chlorine into dilute solutions of calcium chloride at 0°C gives feathery crystals with the composition C12,7-3H20 and with a melting point above 0°C. The crystal is an example of an ice-like hydrate or clathrate containing polyhedra of hydrogen-bonded water molecules linked in the form characteristic of the ice-I structure. There are 46 water molecules in the unit cell which contains 2 small and 6 rather larger holes. If all the voids are filled, as is probably the case wjth argon and small molecules like H 2 S, this corresponds to the formula Μ,5·75Η 2 0, whereas in chlorine hydrate all the larger holes but only about 20% of the smaller holes are occupied250. Bromine forms a similar hydrate with a stoichiometry which is rather ill-defined but which approximates to Br 2 ,8H 2 0, melting at about 6°C. Again X-ray diffraction suggests a relatively open ice-like structure with 172 water molecules in the unit cell of the host lattice and a maximum of 20 polyhedral cavities large enough to accommodate Br2 molecules; occupancy of all these cavities corresponds to a theoretical composition Br 2 ,8-6H 2 0 251 . In some cases of dramatic proportions, the changes in the solubility of the molecular halogens with the addition of halide ions to an aqueous medium arise from the formation of polyhalide ions such as X 3 ~ (X = Cl, Br or I), I2X - (X = Cl or Br) or Br2Cl -. The character­ istics of these species are treated more amply in the context of Section 4 (see p. 1534). Electrophilic Attack of Organic π-Electron Systems204*252 Though as yet unproved, the intermediacy of outer and inner complexes is probably essential to the electrophilic action of the molecular halogens (i) on aromatic compounds 246 j . Sigalla, / . chim. Phys. 58 (1961) 602. 247 H . Davy, Phil. Trans. 101 (1811) 155. 248 M . Faraday, Quart. J. Science, 15 (1823) 71. 249 K. W. Allen, / . Chem. Soc. (1959) 4131. 250 A. F. Wells, Structural Inorganic Chemistry, 3rd edn., p. 578. Clarendon Press, Oxford (1962); H. M. Powell, C. K. Prout and S. C. Wallwork, Ann. Rep. Chem. Soc. 60 (1963) 576; G. A. Jeffrey, Dechema Monograph, 47 (1962) 849; S. Sh. Byk and V. I. Fomina, Russ. Chem. Rev. 37 (1968) 469. 251 K. W. Allen and G. A. Jeffrey, / . Chem. Phys. 38 (1963) 2304. 252 p . B . D . D e la Mare and J. H. Ridd, Aromatic Substitution: Nitration and Halogenation, Butterworths, London (1959); R. O. C. Norman and R. Taylor, Electrophilic Substitution in Benzenoid Compounds, Elsevier (1965); C. K. Ingold, Structure and Mechanism in Organic Chemistry, 2nd edn., Bell, London (1969); P. Sykes, A Guidebook to Mechanism in Organic Chemistry, 3rd edn., Longmans, London (1970).

1218

CHLORINE, BROMINE, IODINE AND ASTATINE: A. J. DOWNS AND C. J. ADAMS

which thus undergo substitution reactions (see example (viii) on p. 1216) or (ii) on olefinic or acetylenic compounds which undergo addition (example (ix)). The formation of the ion-pair which constitutes the inner complex is assisted by a polar medium, e.g. water; furthermore, a second acceptor species A may function as a catalyst or so-called "halogen carrier", presumably through the formation of a transition state such as

+ Α χ

- A = A1X3, ZnX2, IX or ΗΧ (X=C1, Br or I)

The catalyst thus assists the action as a "leaving group" of the halogen atom remote from the aromatic nucleus. Under different conditions supplied, for example, by a non-polar medium, relatively high temperatures, the presence of a radical-initiator instead of a halogencarrier, or irradiation, the interaction may assume a totally different mechanism dominated by homolytic rather than heterolytic fission of the halogen-halogen bond. Thus conditions favouring electrophilic attack lead to substitution at the benzene nucleus rather than addition, which is induced by conditions favouring the formation of halogen atoms. Detailed studies of the electrophilic halogenation of aromatic molecules by molecular chlorine or bromine are consistent with the reaction sequence

X - Cl or Br

the slow step being the formation of a σ-bond between the halogen and the aromatic ring. Molecular iodine is usually unreactive in these circumstances. However, halogenation, including iodination, of aromatic systems can also be achieved by the action of hypohalous acid or a related species (e.g. hypohalite ion or t-butyl hypochlorite) and of interhalogen compounds. Mechanistically the reactions are believed to resemble those involving the molecular halogens. Despite much debate, it seems unlikely that the .electrophilic halogen centre is a cation X + pre-formed by the heterolytic fission of a halogen molecule; a more plausible view is that the incipient cation is conveyed to the aromatic nucleus in the form of a unit such as [XOH 2 ] + , X2, XY, HOX or even perhaps X 3 - (Y = halogen, acetate or alkoxy group). Olefinic or acetylenic units are readily halogenated by addition of chlorine, bromine or interhalogen compounds. Such reactions are of importance in synthesis, while addition of bromine affords a test of unsaturation, in many cases sufficiently quantitative to be useful for analytical purposes. Addition of iodine is less easily accomplished because the thermodynamic balance is much less favourable, the net difference of bond energies providing little advantage to the forward reaction; in general the reaction also proceeds more slowly. Provided the conditions do not favour the formation of halogen atoms, the addition appears to involve a mechanism analogous to that of electrophilic substitution of aromatic nuclei. The following evidence bears witness to such a mechanism: (a) the addition to olefinic systems is normally found to be stereochemically trans and not eis; (b) carrying out the

CHEMICAL PROPERTIES OF THE HALOGENS

1219

addition in the presence of a nucleophile, e.g. halide, N 0 3 - or H 2 0, results in the formation, in addition to the expected dihalide, of products in which the added nucleophile has become bonded to one of the formerly unsaturated carbon atoms; (c) the addition of a "halogen carrier" accelerates the reaction. Consistent with these observations is the mechanism

Br

BrS-

r

Br G

Br |

Br<3

Br

The halogenation of the α-carbon atom of aliphatic aldehydes, ketones, carboxylic acids and related compounds having some capacity to undergo enolisation probably involves a similar sequence. Addition is facilitated by electron-donating substituents attached to the un­ saturated carbon atoms but retarded by electron-withdrawing groups. In extreme cases substituents such as /?-Me2NC6H4- or /?-MeOC6H4-, which have a powerful fortifying action on the donor properties of ethylene, are reported253 to favour the formation of inner complexes incorporating relatively stable organic cations: e.g. Outer complex

Inner complex

R 2 C = C H 2 , B r 2 + Br2

*~

R 2 C=CR 2 . B r 2 + 2Br2

+~

[R 2 C-CH 2 Br] + Br3~"| > R ^ p-Me 2 NC 6 H 4 [R^C-CRJ^B^ j

Formation of Halogen Cations First discussed by Noyes and Stieglitz254 and implicit in many accounts of electrophilic halogenation reactions, the formation of halogen cations of the type X + has been reviewed more critically in recent years255. However, despite the controversy which still surrounds the existence and significance of species such as H2OX + , there is now good evidence of the stable existence of cations such as [py 2 I] + and [(y-pic)2I]+ (py -= pyridine; y-pic=y-picoline); once again these may be viewed as the direct outcome of the transformation of an outer into an inner complex. With a polarizable donor partner, e.g. arsenic or tellurium, such a transformation leads characteristically to oxidation of the donor atom (see reactions (iv) and (vii), pp. 1215-6), but with a strong donor resistant to oxidation, the transformation, though formally similar, corresponds to disproportionation of the halogen (reaction (i)). In either case the change may be assisted by the auxiliary action of a second halogen molecule in forming a polyhalide species with the halide ion displaced from the outer complex. The cations [py 2 X] + , [(y-pic)2X]+ and [(thiourea)2I] + (X = Br or I) produced in this way are characterized by linear coordination about the halogen atom, a feature substantiated by their vibrational spectra215»256 and by crystallographic analysis of the complexes py,2I 2 222 and [(thiourea)2I] + I ~ 257 . Such cations have also been isolated as salts of the nitrate, perchlorate, tetrafluoroborate or hexafluorophosphate anions. 253 R . Wizinger, Chimia, 7 (1953) 273. 254 w . A. Noyes and A. C. Lyon, / . Amer. Chem. Soc. 23 (1901) 460; J. Stieglitz, ibid. p. 797. 255 R . p. Bell and E. Gelles, / . Chem. Soc. (1951) 2734; J. Arotsky and M. C. R. Symons, Quart. Rev. Chem. Soc. 16 (1962) 282. 2561. Haque and J. L. Wood, / . Mol. Structure, 2 (1968) 217. 257 H . Hope and G. H.-Y. Lin, Chem. Comm. (1970) 169.

1220

CHLORINE, BROMINE, IODINE AND ASTATINE: A. J. DOWNS AND C. J. ADAMS

Organic donor molecules which are susceptible to oxidation may give rise to radical-ions as the result of the charge-transfer reaction. Such donors include the polyaromatic compounds perylene, pyrene and carotene and the nitrogen donors /7-phenylenediamine258, Ν,Ν'diphenyl-/?-phenylenediamine186, pyridazine259 and tetramethylhydrazine (example (iii), p. 1215)242. Likewise triphenylamine in propylene carbonate solution is oxidized by iodine to the [/?-Ph2N-C6H4-C6H4-NPh2-/?]+ radical-cation260. In cases such as these the distinction between a redox reaction and formation of a charge-transfer complex largely disappears. Strongly acidic media like sulphuric acid, oleum, fluorosulphonic acid and antimony pentafluoride, individually or in admixture, may favour the formation of cationic species such as X2 + and X 3 + (X = Br or I). For example, when iodine is dissolved in oleum, it gives a deep blue solution containing as a primary ingredient the paramagnetic ion I 2 + 261. For a review of these and other halogen-containing cations, however, the reader is directed to Section 4 (p. 1340). Summary of the Reactions of the Halogens Chemically chlorine, bromine and iodine are well known to be extremely reactive, though the general reactivity is inferior to that of fluorine and decreases in the sequence CI2 > Br2 > I 2 . With a given reagent the halogens commonly give analogous reaction products. That numerous differences do arise, however, is illustrated by the observations (i) that the ultimate products of halogenation of iron are FeCl3, FeBr3 and Fel 2 , of copper are CuCl2, CuBr2 and Cul, and of rhenium are ReCl6, ReBr5 and Rel 4 , and (ii) that although chlorine combines with sulphur dioxide, carbon monoxide and nitric oxide to give sulphuryl, carbonyl and nitrosyl chloride respectively, iodine fails to react in this way. The variations in reactivity and in the products of reactions are determined in part by thermo­ dynamic properties, in part by less clearly defined kinetic factors. Since the enthalpies, and hence free energies, of formation of gaseous halogen atoms or anions do not differ widely for chlorine, bromine and iodine (Tables 1 and 10), the crucial properties which vary considerably and so determine the thermodynamic characteristics of halogenation reactions are (i) the bond energies of molecular halogen compounds or (ii) the sizes of the halide ions through their influence on lattice and solvation energies (see Section 1, pp. 1117-20). With respect to a less electronegative element the halogens appear invariably to follow the bond-energy sequence Cl > Br > I; combination of the halogen with a more electronegative element (oxygen or fluorine) probably leads to a reversal of this sequence. In the reaction with a less electronegative unit M to form a molecular compound MXn, it follows that iodine is at an energetic disadvantage compared with chlorine because of the weakness of the M-I as compared with the M-Cl bond. The disadvantage may be such that Mln is unstable with respect to dissociation into molecular iodine and either M or a lowervalent iodide of M. Conversely iodine forms relatively strong bonds with oxygen or fluorine and, compared with chlorine or bromine, it is thus better disposed to give a compound such as IF m . In these terms it is possible to rationalize, if not to explain, the apparent failure of chlorine or bromine to form a fluoride analogous to IF 7 . On the other hand, if consecutive products of halogenation MX» and MXn +1 can be meaningfully described by the ionic model, variations of stability are primarily a function of the lattice energies of the 258 D . Bargeman and J. Kommandeur, / . Chem. Phys. 4 9 (1968) 4069. 259 R . j . Hoare and J. M . Pratt, Chem. Comm. (1969) 1320. 260 w . H . Bruning, R. F . N e l s o n , L. S. Marcoux and R. N . A d a m s , / . Phys. Chem. 71 (1967) 3055. 261 R. J. Gillespie and J. B . Milne, Chem. Comm. (1966) 158; Inorg. Chem. 5 (1966) 1236, 1577.

1221 aggregates. For a given oxidation state of M, the stability of the compound with respect to the elements then decreases in the order Cl > Br > I, and the energetic advantage of the oxidation MXn -> MX» +x diminishes in the same order (see Section 1, p. 1119). For a given metal of the first transition or lanthanide series, the susceptibilities of the trihalides to the decomposition MX3 ->MX2+£X2 comply with this pattern: accordingly, for M = Eu, Sm or Yb, the decomposition occurs most readily when X = I, while for M = Fe no triiodide is known. Similarly the increase in lattice energy of CuX2 relative to CuX is sufficient to compensate for the (large) second ionization potential of copper in the cases where X = F, Cl or Br but not where X = I. Kinetically two mechanisms are open to such reactions; these may be formally repre­ sented as: CHEMICAL PROPERTIES OF THE HALOGENS

(i) Homolytic

fission

X2

hv, radical Initiator

-> X · + X ·

or Δ

(ii) Heterolytic

fission

X2

Polar solvent

> X+ + X~

Halogen carrier

Homolyticfissioncorresponds to the production of halogen atoms, and the characteristics of reactions in which this mechanism is important have already been discussed (see p. 1165). Reactions in the gas phase or in solution in non-polar solvents generally proceed in this way; the reactions of halogen molecules with hydrogen or hydrocarbons or among themselves come within this category. The foregoing section on complexes of the halogen molecules has provided ample illustration of heterolytic fission, which tends to be the mode of action effective in polar solvents or in the presence of halogen carriers. Some important general reactions of the halogens are summarized in Table 17. The noble gases and nitrogen apart, the halogens react more or less readily with virtually all the other elements. There are exceptions to this generalization: (i) under the action of a microwave discharge chlorine and xenon react to produce xenon dichloride262; (ii) reactions of the halogens with oxygen normally demand forcing conditions in which free atoms and radicals are implicated, though iodine and oxygen are reported to combine directly at high pressure and moderate temperatures32; (iii) although interaction of iodine with sulphur or selenium undoubtedly occurs, definite iodides have not as yet been described. Otherwise the readiness with which reaction occurs depends upon the temperature and phase, upon the state of subdivision, density, rigidity, volatility and solubility of any solid materials partici­ pating in the reaction, and upon the presence of impurities. Reactions of the halogens with solid elements tend to be slow if the products form coherent surface films which are involatile (in a gas-solid reaction) or insoluble (in a solid-liquid reaction). Depending on whether the solid reactant is in a massive or finely divided form, and on whether the surface area is large or small, the rate of reaction may vary dramatically. For example, whereas finely divided aluminium, antimony or copper burns spontaneously in an atmosphere of bromine, such active metals as sodium and magnesium in the massive state do not react with dry bromine even at temperatures up to 300°C; indeed a magnesium vessel has been used for precise measurements of the heat capacity and vapour pressure of bromine. On the other hand, all the halogens interact readily with metal vapours: with alkali-metal vapours, for example, a luminousflameresults and the mechanisms of the accompanying reactions have been the subject of numerous studies32»61. Combustion of certain metals (and compounds) in an atmosphere of chlorine lends itself to calorimetric measurements; hence enthalpies of 262 L . Y. Nelson and G. C. Pimentel, Inorg. Chem. 6 (1967) 1758.

1222

CHLORINE, BROMINE, IODINE AND ASTATINE! A. J. DOWNS AND C. J. ADAMS TABLE 17. SOME GENERAL REACTIONS OF THE HALOGENS 32 » 33

Reaction

Reagent

Comments

1. Donor molecules

Complex-formation with nitrogen-, oxygen-, See pp. 1196-1220. sulphur-, selenium- or halide-donors, or with 7r-donors

2. Metals

2M+«X 2 -*2MX n

With most metals

3. Hydrogen 4. Molecules containing C-H bonds

H 2 +X 2 -*2HX

Free energy of reaction de­ creases in the series Cl >Br > I in parallel with the decreas­ ing bond energies of HX or RX (see Fig. 10) Products and rates depend markedly upon proportions of reactants, phase, tem­ perature, presence of cata­ lyst The relatively low N-X bond energy causes N-X com­ pounds to be strong oxidizing or halogenating agents

R-H+X 2 ->R-X+HX (R = organic group)

5. Hydrides of B, Si or Ge Substitution, e.g. SiH 4 +X 2 ->SiH 3 X+HX B 2 H 6 +6X 2 -> 2BX3 + 6HX 6. Compounds containing (i) Substitution, e.g. CH2CO CH2CO ^N-H bonds

1 \

1 /

7. Unsaturated organic molecules

1 \

NH+Br 2 ->

NBr+HBr

1 /

CH2CO CH2CO RCONH 2 +Br 2 -► RCONHBr+HBr Rearrangement and I hydrolysis j RNH 2 +C0 2 +HBr (ii) Oxidation, e.g. N 2 H 4 +2X 2 -*N 2 +4HX Addition, e.g. X

^c=c'+x 2 -> —c—c—

Hofmann reaction for the synthesis of amines .N-X compounds are pre­ sumably intermediates Free energy of reaction de­ creases in the series Cl > Br > I in parallel with the decreasing C-X bond energy

X 8. Group V element

(i)2M+3X 2 ->2MX 3 (ii)MX 3 +X 2 ^MX 5

9. Group VI element

(i)2E+X 2 -*E 2 X 2 (ii)E+X 2 ->EX 2 (iii)E+2X 2 -*EX 4

10. Other halogens

Formation of interha ogens: X 2 +Y 2 ->2XY X 2 +3Y 2 -*2XY 3 X 2 +5F 2 -*2XF 5 I 2 +7F 2 -*2IF 7

M = P,As,SborBi;X = Cl, Brorl M = P, X = Cl or Br; M = Sb, X = Cl E = S or Se, X = Cl or Br; E = Te, X = I E = S , X = C l ; E = Se,Teor Po, X = Cl or Br E = S, X = Cl; E = Se, X = Cl or Br; E = Te or Po, X = Cl, Br or I X, Y = different halogens Y = F, X = Cl, Br or I; Y = Cl, X = Br or I 1 X = Cl, Br or I

CHEMICAL PROPERTIES OF THE HALOGENS

1223

Table 17 (cont.) Reaction

Reagent

Comments

11. Halide ions

(i) Formation of polyhalide ions, e.g. X 2 + Y - -*X 2 Y(ii) Oxidation: X2 + 2Y-->2X-+Y 2

12. Water

(i) Hydrolysis: X 2 + H 2 0 -> H + + X - +HOX (ii) Oxidation: X2 + H 2 0 - > 2 H + + 2 X - + £ 0 2

Reaction depends on pH. HOX subject to disproportionation

13. Metal oxides

Typically, yX2+2MOe+2zC

With many metal oxides at elevated temperatures; car­ bon not always necessary

14. Metal carbonyls and related compounds

Substitution, e.g. Mn2(CO)io+X2 -> 2Mn(CO)5X Ni(CO) 4 +X 2 -> NiX 2 +4CO

-> 2MXy+2zCO

X, Y = same or different halo­ gen X = Cl,Y = B r o r I ; X = Br, Y= I

Reaction encouraged by photol­ ysis

X = Cl,BrorI

formation have been directly derived, for example, for VC1 4 263 , NbCl5 2<54, TaCl 5 264 , Z r C l 4 264 a n d H f C l 4 263,264.

In their normal (massive) states the following metals are notable for their resistance to attack by the dry halogens at temperatures up to at least 100°C: with respect to gaseous chlorine: nickel, Inconel, Hastelloy, magnesium, Monel, stainless steel, copper, mild steel, cast iron and tantalum; with respect to liquid or gaseous bromine: nickel, magnesium, platinum, lead, tantalum, mild steel and cast iron; with respect to gaseous iodine: magne­ sium, lead, platinum, gold and bismuth. However, few of these metals offer the same resist­ ance in the presence of moisture, which exacerbates the corrosive properties of the halogens. The pronounced catalytic effect of moisture is probably due in part to hydrolysis of the halogen with the formation of the hydrogen halide and hypohalous acid and in part to the provision of a medium for localized electrochemical action on the surface of the metal. The reactions with hydrogen range in their thermodynamic and kinetic properties from the explosive, decidedly exothermic reaction of chlorine to the slow, reversible and barely exothermic reaction of iodine (see Table 29). In that all these reactions are now believed to involve the intermediate agency of atoms, certain features have already been discussed in connection with the formation and properties of the halogen atoms. The kinetics of the reaction between gaseous iodine and hydrogen have been studied in great detail, notably in the classical investigations initiated by Bodenstein32'43: the reaction is thus found to be bimolecular with an activation energy of ca. + 39 kcal mol _ 1 . For many years the reaction was taken to be a model of a bimolecular reaction with no more than a single elementary step, namely the formation of the transition state H2l2, wherein the bonds of the molecules are weakened while incipient bonds corresponding to the products are beginning to form. Such a mechanism has been widely discussed, elaborated and justified by theoretical analysis43»265*. More recently, however, evidence of an atomic mechanism has been found 263 p . Gross and C. Hayman, Trans. Faraday Soc. 60 (1964) 45. 264 L. A. Reznitskii, Zhur. fiz. Khim. 41 (1967) 1482; G. L. Gal'chenko, D. A. Gedakyan and B. I. Timofeev, Russ. J. Inorg. Chem. 13 (1968) 159. 265 (a) K. J. Laidler, Chemical Kinetics, 2nd edn., McGraw-Hill, New York (1965);S. W. Benson, The Foundations of Chemical Kinetics, McGraw-Hill, New York (1960); G. C. Fettis and J. H. Knox, Progress in Reaction Kinetics, Vol. 2 (ed. G. Porter), p. 1, Pergamon (1964); B. A. Thrush, Progress in Reaction Kinetics, Vol. 3 (ed. G. Porter), p. 63, Pergamon (1965); (b) J. H. Sullivan, / . Chem. Phys. 46 (1967) 73.

1224

CHLORINE, BROMINE, IODINE AND ASTATINE: A. J. DOWNS AND C. J. ADAMS

on the basis of studies of the photochemical reaction between hydrogen and iodine265*. Formally similar mechanisms are thus implied for the interaction of hydrogen with chlorine, bromine and iodine32»43'265: hv. A

Initiation:

X2

*X+X

(1)

or wall of vessel

Propagation: X + H 2

> HX 4- H

(2)

H+X2

^HX+X

(3)

H+HX

^H2+X

(4)

Termination: X + X

> X2

(5)

H+X

>UX

(6)

H+H

*H2

(7)

Differences arise because of the increasingly endothermic character of reaction (2) in the series Cl, Br, I (Fig. 10). Whereas reaction chains are set up in the hydrogen/chlorine reaction by repetition of processes (2) and (3), such chains are markedly abbreviated in the hydrogen/bromine and hydrogen/iodine reactions at all but the highest temperatures. First proposed in effect by Nernst266, to account for the photochemical reaction of hydrogen with chlorine, the mechanism implicit in reactions (l)-(7) receives support from numerous findings: for example, (i) the quantum efficiency of the photochemically induced reaction of hydrogen with chlorine is typically 105-106 (cf. 0-01 for the hydrogen/bromine reaction at room temperature); (ii) assumption of a steady-state condition leads to the rate equations d

„, / * i \ * rTTT., -[HBr]therm=2Ä;2(^ -

— [HBr] phot = 2 ·—■ dt (k5)i

[H 2 ][Br 2 ]i

2 **"* l+& 4 [HBr]/fc 3 [Br 2 ]

(/ a b s = number of quanta absorbed)

for the thermal and photochemical reactions between hydrogen and bromine; such expres­ sions are identical in their concentration-dependence with those of experimental measure­ ments32, while the variation with temperature of the term k2(ki/k5)^ gives an activation energy in satisfactory agreement with the algebraic sum +ΔΗ2χ + %ΔΗιχ — $ΔΗ5*; (iii) the presence of free atoms has been detected experimentally. The initiation of the reaction by dissociation of the halogen molecule can be brought about by the action of heat, light or other agents. In the thermal combination of hydrogen with chlorine, dissociation of the chlorine molecules occurs on the walls of the containing vessel, whereas photochemical combination is initiated in the gas phase. In the investigation of these reactions photochemical action has proved important in that it isolates a particular elementary reaction, here the dissociation of the halogen molecule, the rate of which is controlled by the light intensity. Hence the importance of that reaction may be judged, and its elimination allows a closer inspection of the other steps. For this and other reasons, the influence of light on the combination of hydrogen with the halogens has been extensively studied; it is probable that the photo­ chemical reaction of hydrogen with chlorine has been more thoroughly examined than any other, ever since the days of Bunsen. For each system the effects of pressure, temperature and frequency of the irradiating light have been explored in detail. Investigations have also been concerned with the influences of the walls of the reaction vessel, the presence of "third bodies" in the gas phase (e.g. inert gases), and the inhibiting effect of agents such as oxygen or nitric oxide which have the capacity to terminate the reaction chains (see p. 1166). In practical terms, the formation reactions of the hydrogen halides have been suggested as the basis of a fuel cell267. 266 w . Nernst, Z. Elektrochem. 2 4 (1918) 335. 267 w . V. Childs, U . S . Pat. 3,445,292 (1968).

CHEMICAL PROPERTIES OF THE HALOGENS

1225

An atomic mechanism is almost certainly responsible for the reactions between the halo­ gens themselves. According to the conditions employed, fluorine reacts (a) with chlorine to form C1F, CIF3 or C1F5, (b) with bromine to form BrF, BrF 3 or BrF5, and (c) with iodine to form IF, IF 3 , IF 5 or IF 7 . The products of the interaction of chlorine with bromine are BrCl or BrCl3 and related compounds268; iodine and chlorine afford IC1 or I2C16; IBr is the only product as yet identified in the reaction between iodine and bromine. Certain reactions require forcing conditions, e.g. Excess fluorine

C1 2 +5F 2

^2C1F5

Total pressure 250 atm, 350°C Gas phase

Br 2 +3C1 2

Microwave discharge

► 2BrCl3

(ref. 268)

but others take place with less inducement, e.g. Room temperature

Br 2 +C1 2 ^

N

2BrCl

Gas phase or solution Room temperature

I 2 + 5F 2 I2+3F2

CC13F suspension. -45°C

► 2IF 5 *2IF3

(ref. 269)

These reactions and the properties of the interhalogen compounds thereby produced are discussed in Section 4 (p. 1478). Neither chlorine nor bromine reacts with molecular oxygen under normal conditions, but at 325°C and with a partial pressure of oxygen of ca. 1200 atm iodine is reported to give a yield of 2-3% of iodine pentoxide32. The action of oxygen atoms on chlorine gives rise to the molecules ClO and C10 2 ; the kinetics of the primary reaction

o+ci2-»cio+ci

have been studied at 300°K by discharge-flow coupled with mass-spectrometric methods of detections^. Likewise ozone and chlorine yield the relatively short-lived CIO3 in a reaction which is susceptible to photolytic action; irradiation by blue light gives detectable amounts of C1207. By the action of atomic oxygen or ozone on bromine under suitable conditions bromine oxides are obtained. Chlorine atoms produced, for example, by photolysis combine with oxygen to give, in the first place, the peroxy free-radical ClOO, which then plays a central role in the formation and subsequent decay of the ClO radical according to the mechanism132,270 : ClO+ClO ^Cl+ClOO

cioo+ci->ci2+o2

ClOO+M x± Cl+0 2 +M (M = third body) As produced by flash photolysis of chlorine-oxygen mixtures, ClO has a half-life of a few milliseconds. Nevertheless, it has been possible to explore the formation and fate of species such as ClO and ClOO by measurements of their electronic absorption spectra; ClO has also been characterized by its esr93 and microwave271 spectra. Formed by the action on bromine vapour of oxygen atoms produced in a microwave discharge, the gaseous BrO molecule has also been characterized by its esr93 a n c i microwave272 spectra, while Br0 2 has been isolated 268 L . Y. Nelson and G. C. Pimentel, Inorg. Chem. 7 (1968) 1695. 269 M . Schmeisser, W. Ludovici, D. Naumann, P. Sartori and E. Scharf, Ber. 101 (1968) 4214. 270 M. A. A. Clyne, Ann. Rep. Chem. Soc. 65A (1968) 173-174, 183-185. 271T. Amano, S. Saito, E. Hirota, Y. Morino, D. R. Johnson and F. X. Powell, / . Mol. Spectroscopyy 30 (1969) 275.

1226

CHLORINE, BROMINE, IODINE AND ASTATINE: A. J. DOWNS AND C. J. ADAMS

at low temperatures following the action of an electric discharge on a mixture of bromine vapour and oxygen32. The intermediate formation of BrO or other bromine oxides probably accounts for the action of bromine as a "sensitizer" for the photochemical decomposition of ozone, ClO or C102. Chlorine reacts with many metal oxides to yield anhydrous chlorides or oxychlorides, the presence of carbon commonly being necessary to expedite the reaction. Bromine and iodine are less prone to react with metal oxides, and, where attack does occur, oxyhalogen species such as bromate or iodate tend to be more prominent as products. Thus, when iodine vapour mixed with oxygen is passed over heated alkaline-earth oxides or carbonates, solid periodates are formed: e.g. 10CaO + 2 I 2 + 7 O 2 -> 2Ca 5 (I0 6 ) 2

With mercuric oxide, however, chlorine gives the oxide C120 in a reaction which affords a standard method of preparation; stable only at temperatures below — 50°C, Br 2 0 is likewise formed, but I 2 0 is not known. Alkali or alkaline-earth metal chlorites are oxidized by chlorine to chlorine dioxide (Table 14). This process, which is important for the industrial production of the dioxide, can be adapted for either continuous or batch operation, and can be carried out under dry or aqueous conditions. Hydrogen peroxide is initially oxidized by chlorine, bromine or iodine according to the equation H 2 0 2 +X2->0 2 +2HX

Hydrogen chloride, once formed, is stable, but oxidation by hydrogen peroxide causes the bromide or iodide to revert to the free halogen. The net result of the interaction with bromine or iodine is therefore the catalyzed decomposition of hydrogen peroxide. The kinetics of these reactions in aqueous solution have been the subject of considerable investi­ gation, on the basis of which possible mechanisms have been discussed32. By the formation of halogen atoms, photolysis of the reaction mixture initiates chain reactions in which propagation steps such as X+H 2 0 2 ->H0 2 4-HX

probably participate. By contrast with the behaviour of hydrogen peroxide, peroxydisulphuryl difluoride FS0 2 OOS0 2 F serves as a source of F S 0 2 0 radicals which oxidize the halogens to produce halogen fluorosulphates. Chlorine is thereby converted to C10S0 2 F, bromine to BrOS0 2 F or Br(OS0 2 F) 3 (depending upon the conditions), and iodine to I 3 OS0 2 F, IOS0 2 F or I(OS0 2 F) 3 (again depending upon the conditions)273. These products are to be regarded as derivatives of the halogens in positive oxidation states (see Section 4). Compounds of the type I(OCOR)3 containing formally electropositive iodine atoms are also formed by the action of fuming nitric acid on iodine in the presence of the appropriate acid anhydride (RCO) 2 0 [R = CH3, CH2C1, CHC12, CC13, CF 3 , C 3 F 7 or C6F5po>274. A similar reaction affords the compound IP0 4 . Interaction of a halogen with a silver(I) or mercury(II) oxysalt in a non-aqueous medium also gives rise to halogen derivatives of the oxyanion: e.g. IC104 ™>™, I(C104)3 ™>™, pyIN0 3 ^ο a n d pylOCORi2o>274 ( p y = pyridine; R = organic group). A compound described as "chlorine acetate" and formu­ lated as C10COCH3 is likewise the outcome of such a reaction, but the structures of this and related compounds have yet to be elucidated. 272 F. X. Powell and D. R. Johnson, / . Chem. Phys. 50 (1969) 4596. 273 A. A. Woolf, New Pathways in Inorganic Chemistry (ed. E. A. V. Ebsworth, A. G. Maddock and A. G. Sharpe), p. 351, Cambridge (1968). 274 M . Schmeisser, K. Dahmen and P. Sartori, Ber. 100 (1967) 1633. 275 N . W. Alcock and T. C. Waddington, J. Chem. Soc. (1962) 2510.

1227 \ Compounds containing ^N-H bonds are susceptible to oxidation by the halogens. For example, under aqueous conditions ammonia and hydrazine are oxidized, at least in part, to nitrogen (see Table 14). However, substitution reactions are also possible, whereby the hydrogen is replaced by a halogen atom. Obtained in this way are compounds such as NH2C1, NCI3, ΝΙ3,ΝΗ3, and N-bromosuccinimide. The ^ N-X bond is characteristically weak with the result that the compounds are powerful oxidizing or halogenating agents; several of them undergo spontaneous decomposition at the slightest provocation, but others, e.g. N-bromosuccinimide, serve as useful halogenating agents for organic systems. The instability of the N-X unit is well illustrated by the Hofmann reaction whereby organic / amides are converted to amines: CHEMICAL PROPERTIES OF THE HALOGENS

O / RC-NH 2 -

O

o ->RC-NHBr-

->[RC-N]-

Rearrangement

->RN=C=0 -> RNH 2

By the action of an elemental halogen X2, M-H, M-M and M-C bonds formed by a metal or metalloid M are commonly supplanted by M-X bonds in the formation of binary or mixed halide derivatives of M. A selection of reactions belonging to this category is depicted in Fig. 18276 ~278. The interaction is occasionally violent and difficult to control, as (*-C 5 H 5 )M(CO) n X e.g.(7r-C5H5) Fe(CO) 2 I or(^-C 5 H 5 )Ni(CO)I

M(CO),Xy, e.g. Fe(CO)4X2 \ or MXZ, e.g. NiX2 v^[ ( 7 t _c 5 H 5 ) M (CO),

XM(CO), (M -Mn, Tc or Re; X=C1, B r o r l ) R,MX (M Si, Ge, Sn or Pb; X -Cl, Br or I)

M2(CO)10

SiH4.nXn(X Cl, B r o r l )

R n _ m MX e.g. M B , Ga, Ge, Sn Pb, Sb, Bi, Pt or Au 1,2- or 1J-C 2 B, 0 H 12

BX 3 (X-C1, B r o r l ) B,HsBr

l,2-orl,7-C 2 B 10 H |2 _ n Cl n (n=2-12) U-orlJ-C.B^Br^l-S)

R ^organic group

B

i2H 12 _ n Xl(n= 1-12) (X-Cl, Br or I)

B f t . f r 1 or 2) e.g. X=Br or I

FIG. 18. Some representative reactions of halogens with M-H, M-C and M-M bonds, where M is a metal or metalloid. 276 E . A. V. Ebsworth, Volatile Silicon Compounds, Pergamon, Oxford (1963). R. M. Adams (ed.), Boron, Metallo-boron Compounds and Boranes, Interscience, New York (1964); M. F. Hawthorne, The Chemistry of Boron and its Compounds (ed. E. L. Muetterties), p. 223. Wiley, New York (1967). 278 G. E. Coates, M. L. H. Green and K. Wade, Organometallic Compounds, 3rd edn., Methuen, London (1967-8). 277

1228

CHLORINE, BROMINE, IODINE AND ASTATINE: A. J. DOWNS AND C. J. ADAMS

between silane and chlorine at room temperature; conversely certain reactions are relatively sluggish at normal temperatures. However, with judiciously chosen conditions of temper­ ature, reaction medium and, where necessary, catalysis, many of the reactions proceed smoothly, affording useful preparative routes, for example, to organometallic halides such as Fe(CO)4I2 or (7r-C5H5)Ni(CO)I. In aqueous solution the azide ion is ultimately oxidized to nitrogen, notably by bromine or iodine, in a reaction which is catalyzed by a number of sulphur-containing species, e.g. H 2 S, CS2, SCN-, S w 0 6 2 ~ (n = 4, 5 or 6), N 3 SCS~and cysteine. Advantage is taken of this catalytic action on the iodine-azide reaction to provide a sensitive test for sulphides and sulphur compounds32»33. The kinetics of the reactions and their likely mechanisms have been discussed33»279,280. Under different conditions the interaction of azides with the molecular halogens has been shown to give the halogen azides XN3 281 . The action of the halogens on other metal pseudohalides may likewise lead to compounds such as XCN, though the course of the reaction is profoundly influenced by the choice of conditions: e.g. (i) C N - + X 2

Aqueous solution or

>XCN+X-

(X = Cl, Br or I)

dry state 0°C, non-aqueous

(ii) 2 M S C N + B r 2

> 2MBr+(SCN>2 solvent, e.sr. CHC13

I TV-y £ 1 .

2C1SCN (M = N H 4 , N a , K or £Pb)282 (iii) SCN " + 4 I 2 + 4 H 2 0

Aqueous soln

> SO4 2 " + I C N + 7 1 " + 8H +

pH7 Aqueous soln

(iv) SCN" + 3 I 2 + 4 H 2 0

PH<7

>S042~

+61" + HCN+7H+

Apart from its importance as a brominating agent in organic chemistry45, bromine is a useful oxidizing agent, though under the aqueous conditions commonly employed it is not always certain whether the active agent is Br2 or a hypobromite species. One facet of these oxidizing properties is presented by the degradation of organic amides (see above), while another consists of the oxidation of aldose sugars to δ-lactones, which are subsequently hydrolysed to aldonic acids. For example, D-glucose can be converted to D-gluconic acid by oxidation with bromine water; the mechanism, which appears to be highly selective, is believed to involve electrophilic attack by bromine on the glycosidic oxygen atom283. Extensive use of bromine oxidation was made in the classical studies of Fischer and others on the structure of sugars. The dissociation of halogen molecules into atoms by the absorption of light is sometimes able to initiate chain reactions between two other substances, the consumption of halogen in the reactions being frequently small or negligible. This property forms the basis of the "sensitizing" action of the halogens on certain reactions. As examples may be cited the chlorine-sensitized combination of carbon monoxide and oxygen, explosion of hydrogen with oxygen, and the oxidation of various organic compounds, the bromine-sensitized decomposition of ozone, and the bromine- or iodine-sensitized decomposition of hydrogen peroxide. It is possible that a similar mechanism operates, at least in

279 j . Weiss, Trans. Faraday Soc. 43 (1947) 119. 280 G. Dodd and R. O. Griffith, Trans. Faraday Soc. 45 (1949) 546. 281 K. Dehnicke, Angew. Chem., Internat. Edn. 6 (1967) 240. 282 A . B. Angus and R. G. R. Bacon, / . Chem. Soc. (1958) 774; R. G. R. Bacon and R. S. Irwin, ibid. p. 778. 283 D . H. Hutson and D . J. Manners, Ann. Rep. Chem. Soc. 61 (1964) 431.

ANALYTICAL DETERMINATION OF THE ELEMENTARY HALOGENS

1229

some cases, in the catalytic action of small quantities of iodine on many organic reactions, including halogenation, metalation, dehydration, isomerization and pyrolytic decomposition. Because of this action, iodine is widely employed to accelerate such reac­ tions or to shorten otherwise excessive induction periods. The role of iodine as a catalyst in the hydrogen-chlorine reaction has been discussed, as has its effect on the ignition temperatures or explosion limits of combustion processes involving hydrogen- or hydro­ carbon-oxygen mixtures32. 2.8. ANALYTICAL DETERMINATION OF THE ELEMENTARY HALOGENS32284.285 Chlorine

The detection and determination of free chlorine depend primarily on its oxidizing action. Thus the release of iodine from iodide has been extensively used both as a qualitative test and as a means of quantitative analysis. The familiar starch-iodide test for chlorine may be used to detect as little as 0· 1 ppm under aqueous conditions. Chlorine in the atmosphere has been detected by its action on potassium iodide deposited on silver leaf32; more commonly the iodine liberated in an aqueous medium is estimated by titration with sodium thiosulphate or arsenite. Another test widely favoured for the detection and estimation of free chlorine involves the use of ö-tolidine (o^'-dimethylbenzidine); in dilute solution in fairly strong hydrochloric acid this gives a yellow coloration, the intensity of which provides a measure of the concentration of chlorine present. The starch-iodide and ö-tolidine tests are thought to be about equally sensitive, but the latter is more specific and so less vulnerable to interference from other substances. Even the o-tolidine test has considerable shortcomings; not only do other free halogens give a positive reaction, but ions such as Mn 3 + , NO2 " and Fe 3 + also cause severe complications. The search for alternatives to this popular test has yielded other colorimetric procedures involving reagents such as 3,3'-dimethylnaphthidine, which yields a purple-red semiquinoid product and is reported to be superior to o-tolidine, /7-aminodimethylaniline, which gives a red coloration, and methyl orange, which is bleached by chlorine. A useful and rather specific test for chlorine (or bromine) depends upon the König reaction with aqueous cyanide ions to produce C1CN (or BrCN), which gives an intensely coloured di-anil derivative with pyridine and an aromatic amine, e.g. benzidine; the method is applicable to chlorine concentrations up to 2 ppm. Other techniques for estimating gaseous chlorine involve absorption in alkali in the presence of a reducing agent, the chloride ions thus formed being determined either volumetrically with silver or mercuric nitrate or gravimetrically with silver nitrate. Chlorine gas is assayed commercially by absorption in mercury, with which it reacts quantitatively, followed by volumetric measurement of the residual gases. Alternatively the chlorine may be absorbed in an acetate-buffered aqueous solution of arsenite. In recent years gaschromatographic methods have been used increasingly for the assay of chlorine gas and the determination of impurities, while infrared spectroscopy has been applied to the detection of organic impurities and moisture. Bromine45

In common with chlorine, free bromine is conveniently estimated by iodometric titration; 284 G. W. Armstrong, H. H. Gill and R. F. Rolf, Treatise on Analytical Chemistry (ed. I. M. Kolthoff, P. J. Elving and E. B. Sandell), Part II, Vol. 7, p. 335. Interscience (1961). 285 F . D. Snell and L. S. Ettre (eds.), Encyclopedia of Industrial Chemical Analysis, Vols. 8, 9 and 14. Interscience (1969-71).

1230

CHLORINE, BROMINE, IODINE AND ASTATINE! A. J. DOWNS AND C. J. ADAMS

bromine vapour can, for example, be quantitatively absorbed in potassium iodide solution and the liberated iodine determined by titration with thiosulphate. Bromine in the vapour phase can also be absorbed in alkali and either (a) reduced, e.g. with the aid of hydrogen peroxide, to the bromide ion, which as AgBr can be estimated gravimetrically or volumetrically, or (b) oxidized, e.g. by hypochlorite, to the bromate ion, which lends itself to iodometric titration. However, the detection of bromine is often accomplished most easily by one of several colour tests. The colour of free bromine, particularly in carbon tetrachloride or carbon disulphide solution, itself serves as a method of detecting and esti­ mating the element. Various organic reagents which form coloured compounds with bromine are also available, though few are specific; thus most of the reagents which give colorations with chlorine respond similarly to bromine. Reagents recommended for the detection and estimation of free bromine include (a) fluorescein or, better, ethylfluorescein, which affords the red dye eosin; (b) dithizone, which gives an intense red colour in carbon tetrachloride solution; (c) rosaniline (fuchsin), which undergoes bromination to give violet bromoderivatives; (d) phenol red, which is converted to an indicator of the bromophenol-blue type. Despite problems of interference by other oxidizing agents, several of these colour tests have been successfully applied to the determination of trace quantities of elementary bromine. However, one of the most sensitive procedures for estimating bromine results from its catalytic action on the oxidation of iodine to iodate by acid permanganate; in one such procedure the change in iodine concentration is determined spectrophotometrically using a carbon tetrachloride extract of the aqueous solution. Various methods have been devised for the evaluation of impurities in liquid bromine. The chlorine content can be determined by measurements of the density of the liquid or of the depression of freezing point, while in amounts of the order 5-10% the halogens can be reduced with sulphur dioxide to the corresponding halides, which are then conveniently estimated by potentiometric titration with silver nitrate. Pure bromine being transparent throughout the normal infrared region, moisture and most organic impurities have been detected with considerable sensitivity by measuring the infrared spectrum of the liquid54. Iodine Free iodine in dilute aqueous solution is readily determined by direct titration with a standard solution of sodium thiosulphate or sodium arsenite, typically using starch as the indicator; the equivalence point may also be determined electrometrically, for example by the so-called "dead-stop" technique. Other methods of volumetric analysis include oxidation by potassium iodate in the presence either of concentrated hydrochloric acid or of cyanide ions, whereby iodine (or iodide) is converted to IC12 ~ or ICN; iodine may also be oxidized to iodate with bromine or permanganate, and the iodate estimated by iodometric titration, six moles of I 2 being thereby released for each mole originally present. Iodine can be detected as the free element by the violet colours of its vapour and of its solutions in solvents such as carbon tetrachloride, chloroform or carbon disulphide. Perhaps the most widely used test is afforded by the blue-black coloration due to the starchiodine reaction, the sensitivity of which has been studied as a function of numerous factors. α-Naphthoflavone has been recommended as a test for free iodine more sensitive than starch, 0-1 ppm being readily detected; malachite green has also been shown to be superior to starch in very dilute solutions or in the presence of large amounts of electrolyte or ethanol286. 286 j . o . Meditsch, Analyt. Chim. Acta, 31 (1964) 286.

1231

BIOLOGICAL ACTION OF THE ELEMENTARY HALOGENS

Colorimetric methods favoured for the determination of small quantities of free iodine involve the starch-iodine complex, the α-naphthoflavone reagent, the I3 - ion formed in aqueous solution, or solutions of iodine in organic solvents. 2.9. B I O L O G I C A L A C T I O N OF THE E L E M E N T A R Y HALOGENSP2.33.45.284.285.287

As vapours, all of the halogens cause similar physiological reactions. Through their oxidizing action, comparatively low concentrations of the vapours are highly irritating to the entire respiratory tract, the mucous membranes and the eyes, producing responses such as coughing and smarting of the eyes; approximate exposure limits are given in Table 18. TABLE 18. EXPOSURE LIMITS FOR THE HALOGEN VAPOURS 46 .284,287

Volume concentrations in ppm

Maximum allowable concentration that can be tolerated without disturbance: Prolonged exposure (several hours) Short exposure (£-1 hr) Least detectable odour Least amount causing immediate irritation to the throat Dangerous to life even on short exposure ($-1 hr) Rapidly fatal on very short exposure (< £hr)

Chlorine

Bromine

Iodine

0-35-1 0 4 3-5 10-15 40-60 1000

01-1 4 -3-5 15 40-60 1000

01 0-3

Exposure to high concentrations of the vapours leads to inflammation and congestion of the respiratory system and oedema of the lungs, which in severe cases can be fatal. Because of their greater volatility, chlorine and bromine are much more likely to be encountered in dangerous concentrations than is iodine. Accordingly, few cases of iodine poisoning have been recorded, though this apparently innocuous record should not disguise the lachry­ matory action of the vapour or the fact that exposure can evoke pulmonary oedema, eye irritation, rhinitis, chronic pharyngitis and catarrhal inflammation of the mouth. By con­ trast, the toxic nature of gaseous chlorine is well recognized, as testified by its history as one of the first gases to be used in chemical warfare in World War I. The odours of the halogens give adequate warning against acutely dangerous exposure but not against concentrations which may be harmful as a result of prolonged exposure. However, there have as yet been no reports of cumulative intoxication following several exposures which do not individually exceed the normal safety limits. Liquid chlorine is a skin irritant causing burns on prolonged contact; it can also cause severe eye damage. Likewise liquid bromine presents a serious hazard in handling. It rapidly attacks the skin and other tissues to produce irritation and necrosis; the burns are not only painful but slow to heal. In short-term contact with the skin, solid iodine causes a discoloration similar to that produced by "tincture of iodine", but prolonged contact can cause burns or allergic dermatitis. The manipulation of elementary chlorine, bromine and iodine invariably demands the following safety measures: (i) rigorous avoidance of all contact of the halogen with skin, 287 M. B. Jacobs, The Analytical Toxicology of Industrial Inorganic Poisons, p. 635. Interscience (1967).

1232

CHLORINE, BROMINE, IODINE AND ASTATINE: A. J. DOWNS AND C. J. ADAMS

eyes and clothing, and (ii) adequate ventilation with all possible precautions against inhaling the vapours. Where significant quantities of chlorine or bromine are to be handled, good practice requires that goggles and gloves are worn and that suitable gas masks are available. Preventive health measures and first-aid procedures are given elsewhere288. The poisonous action of the free halogens on micro-organisms affords an important means of disinfection. Thus in the treatment of drinking water, the use of small quantities of chlorine has become standard practice, and by using higher halogen concentrations the bacterial count of effluents from industrial plants can be controlled. The activity of chlorine in the destruction of organisms more complex than bacteria, e.g. non-pathogens such as Crenothrix formed in pipe-lines, has also been investigated. Likewise dilute aqueous solutions of bromine exert considerable germicidal action; comparative studies suggest that it is commonly intermediate between chlorine and iodine in its activity on bacterial spores, though with respect to some species, e.g. yeast and certain algae, it is said to be more toxic than chlorine. For many years iodine has been used in various forms as a first-aid antiseptic for the treatment of cuts and abrasions, and as routine surgical procedure in the sterilization of operation sites; despite the advent of many antiseptics of higher activity, the use of iodine for these purposes has persisted. The specific action of iodine solutions on a variety of organisms has been widely investigated; the solutions have thus been shown to exhibit both bacteriostatic and bactericidal activity. The halogens are known to react with many materials found in living matter, e.g. unsaturated aliphatic acids (by addition), aromatic amino-acids (by ring-substitution), amino-derivatives (with oxidation or formation of N-halogen compounds), organic donor species (to form molecular complexes which may, for example, be relatively insoluble in water), and thiols and disulphides (with oxidation). While it is not certain that all of these reactions take place in vivo, it is conceivable that any one of them could contribute to the biological effects peculiar to the elementary halogens. In specific cases the toxic action of the halogens may well depend upon oxidation of enzyme functional groups such as -SH and -SS-; it is now fairly conclusively proved, for example, that the activity of chlorine towards various micro-organisms is by complete inhibition of the sulphydryl enzymes of bacteria, leading to their death.

3. HALIDE I O N S AND RELATED SPECIES: OXIDATION S T A T E - 1 3.1. PROPERTIES OF THE HALIDE IONS

Thermodynamic Properties, Ionic Radii and Polarizabilities Reduction of a free halogen atom leads to the formation of a halide ion with the closedshell configuration ns2np6 and electronic ground state lS; for all the halogens this is a thermodynamically spontaneous process. To the extent that the simple ionic model is applicable to real chemical systems, the halide ions are recognizable entities, to which can be attributed sizes as well as spectroscopic, thermodynamic and magnetic properties. However, even in the most favourable circumstances the simple ionic model, with its impli­ cations of a hard sphere bearing unit charge, is unrealistic; in all environments of chemical significance some transfer of charge between the ion and its neighbours is inevitable. Thus, although Table 19 lists some of the better defined physical characteristics of the chloride, 288 See for example Chlorine Manual, 3rd edn., The Chlorine Institute, New York (1959).

PROPERTIES OF THE HALIDE IONS

1233

TABLE 19. PHYSICAL PROPERTIES OF THE HALIDE IONS

cr

Property Electron configuration and ground state Ionic radius (6 : 6-coordination) (A) Ionization potential (elec­ tron affinity of halogen atom) at 298°K Thermodynamic properties at 298°K (i) of gaseous anion: -A#,°(kcal) 5°(caldeg-imol-i) (ii) of the hydrated anion -AJV(kcal) -AG / °(kcal) S°(caldeg-imol-i) Spectroscopic properties X-ray spectra X-ray scattering factors Ultraviolet spectra in aqueous solution: wavelength (A) [frequency (cm ~ *)]; oscillator strength (f) Electrical and magnetic pro­ perties Polarizability (A 3 ) Diamagnetic susceptibil­ ity ( x l 0 6 c.g.s. units per g ion) Standard potential for the aqueous system *Χ2/χ-,£°(ν) Ionic mobility (limiting equivalent conduc­ tance) in water at 298°K (Λ cm 2 o h m - 1 g ion - 1 ) a

[Ne]3j23/>6

15

I"

[Ar]3i/i04524^6 15

[Ki]4du>5s25p* 15

1-96

219

1-81 eV 3-68

kcal 84-8

Ref.

Br"

eV 3-43

kcal 790

eV 3-13

a kcal 72-1

a a,b

52-3 46-6 55-7 391 40-4 36-7 rel. ÖH+ = 1 rel. gas­ rel. ΛΗ+ = 1 rel. gas­ rel. ÖH+ = 1 rel. gas­ eous ion eous ion eous ion 39-933 88 29039 80 13-60 70 31-383 82 24-900 75 12-44 67 13-57 20 19-91 16 25-60 11 Refs. c-e Ref. g

1820 [55,000]; 0 0 9

3-475

Refs. d, e Ref. g

1905 [52,500]. n 9 R 2020 [49,500]' U Z Ö

4-821

Refs. d, e Ref. g

1940 [51,500], 0 , 4 7 2260 [44,300]'

7-216

f

h

-23-4

-34-6

-50-6

i

+ 1-356

+ 1065

+ 0-535

a

76-35

78-14

76-8 4

j

A. G. Sharpe, Halogen Chemistry (ed. V. Gutmann), Vol. 1, p. 1, Academic Press (1967). National Bureau of Standards Technical Note 270-3, U.S. Govt. Printing Ofl&ce, Washington (1968); Codata Bulletin, International Council of Scientific Unions, Committee on Data for Science and Technology (1970). c Gmelins Handbuch der Anorganischen Chemie, 8 Auflage, "Chlor", Teil A, p. 133 (1968). a Mellor's Comprehensive Treatise on Inorganic and Theoretical Chemistry, Supplement II, Part I, pp. 624-626, 776-778, 910-912, Longmans, Green and Co., London (1956). e J . A. Bearden, X-ray Wavelengths, U.S. Atomic Energy Commission, NYO-10586, Oak Ridge, Tennessee (1964); J. A. Bearden, Rev.Mod. Phys. 39 (1967) 78; J. A. Bearden and A. F. Burr, ibid. p. 125; A. E. Sandström, Experimental Methods of X-ray Spectroscopy: Ordinary Wavelengths, Handbuch der Physik, 30 (1957) 78. * J. Jortner and A. Treinin, Trans. Faraday Soc. 58 (1962) 1503. «International Tables for X-ray Crystallography, Vol. Ill (general ed. K. Lonsdale), p. 201, Inter­ national Union of Crystallography (1962); H. P. Hanson, F. Herman, J. D. Lea and S. Skillman, Acta Cryst. 17 (1964) 1040; D. T. Cromer and J. T. Waber, Acta Cryst. 18 (1965) 104. h K. Fajans, / . Phys. Chem. 74 (1970) 3407. 1 A. Earnshaw, Introduction to Magnetochemistry, p. 5, Academic Press, London (1968). i R. A. Robinson and R. H. Stokes, Electrolyte Solutions, 2nd edn., p. 463, Butterworths, London (1959). b

1234

CHLORINE, BROMINE, IODINE AND ASTATINE! A. J. DOWNS AND C. J. ADAMS

bromide and iodide ions, little absolute significance can be attached to those values, e.g. of ionic radii, which are derived from measurements focused on the halide ion in crystalline solids or in solution. The importance of such results resides not in their accuracy of des­ cription but in the success with which they can be applied, through the medium of the ionic model, to the interpretation of the thermodynamic and crystallographic properties of halide systems. The enthalpy of formation of a gaseous halide anion, — Ä///[X~(g)], decreases in the sequence of Cl~ > Br _ > I~. The standard entropy of the gaseous anion is lower than that of the corresponding atom because of the difference in electronic multiplicity; according to the Sackur-Tetrode equation, the contribution to the standard entropy of a monatomic gas of electronic multiplicity Qe is R In Qe = 0 or 2-75 cal deg - 1 for a gaseous halide ion or halogen atom, respectively289. Since the standard entropies and entropy changes vary but little from halogen to halogen, variations in — AG/[X_(g)] follow the sequence dictated by -AHf[X-(g)]. The fact that halides of pre-transition metals have internuclear distances which are approximately additive suggests the feasibility of assigning radii to the individual ions290. The radii adopted in Table 19, which refer to a sodium chloride structure, are taken from Sharpe's review of the physical inorganic chemistry of the halogens289. In this context, however, two complications are noteworthy, (i) For the very few crystals in which the variation in electron density along a line joining the nuclei has been determined, the minimum in electron density does not occur at the point indicated by the generally accepted ionic radii. In NaCl, for example291, the minimum occurs at 1-64 A from the chlorine nucleus. (ii) The ionic radius varies somewhat with the coordination number of the halide ion. This is a consequence of the dependence of the lattice energy of an ionic halide on the Madelung constant A of the lattice-type, viz.

where z+ is the charge on the cation, e the electronic charge, N the Avogadro number, r the equilibrium interionic separation and p a function of the compressibility (almost constant), which takes into account the interionic repulsion arising from the finite size of the ions. As the number of cation-anion contacts increases, not only A but also r becomes larger. As a result of these mutually compensating changes, there appears to be little difference in energy between different structures, e.g. CsCl and NaCl types, in .justification of the empirical formulae UL =

256vz+

r++r-

or

287vz+ r

r++r- i

1

0-345 i

r++r-]

kcalmol - 1

(2)

(v = number of ions per mole of halide; r+ and r_ = radii of cation and halide anion respectively) developed by Kapustinskii and Yatsimirskii292 for the treatment of ionic radii and lattice energies in the absence of structural details. Approximate though the method is, its simplicity makes it highly suitable for comparative purposes, and the conclusions ob­ tained by its use often hold even when the compounds involved are far from ionic289»293. 289 A. G. Sharpe, Halogen Chemistry (ed. V. Gutmann), Vol. 1, p. 1. Academic Press (1967). 290 L . Pauling, The Nature of the Chemical Bond, 3rd edn., Cornell University Press, Ithaca (1960). 291 H . Witte and E . Wölfel, Z. phys. Chem. (Frankfurt), 3 (1955) 296. 292 A . F . Kapustinskii, Quart, Rev. Chem. Soc. 10 (1956) 2 8 3 ; T. C. Waddington, Adv. Inorg. Chem. Radiochem. 1 (1959) 157. 293 D . A . Johnson, Some Thermodynamic Aspects of Inorganic Chemistry, Cambridge (1968).

PROPERTIES OF THE HALIDE IONS

1235

Hence the radius of a single ion has no precise physical basis, the chief justification for values such as those given in the table lying in their practical usefulness in comparative interpretation and in their capacity roughly to reproduce internuclear distances and account for those properties to which such distances are related. The enhanced susceptibility to charge-delocalization which accompanies the increasing size and nuclear charge of the halide ions is indicated (i) by the polarizabilities and (ii) by the ionization potentials of the ions. Correspondingly the donor capacity of the ions decreases in the order I - > Br~ > Cl~. Primary interactions of the type M n + · · · X~ therefore involve a proportion of electron-delocalization (that is, charge-transfer or ulti­ mately covalent bonding), which is greater for X = I than for X = Cl; secondary inter­ actions, e.g. dispersion forces, are also relatively stronger in systems containing the larger halide ions. For this reason, structures which give a good approximation to the ionic model are more commonly formed by fluoride than by iodide ions, the larger size and polarizability of which tend to favour lower coordination numbers, often in conjunction with layer or chain structures. With cations of high electron affinity, formed by non-metals or by metals in high oxidation states, the halide ion cannot survive as a recognizable entity, charge being transferred roughly in accordance with Pauling's Electroneutrality Principle, which requires that, for stability, the charge on any atom of an aggregate shall not greatly exceed ± Je 294 . Such systems may take the form of neutral molecules, e.g. SiX4 or WC16, or of complex ions, e.g. HgX 4 2 ~ or PdX 6 2 ~. Donor Properties of the Halide Ions: Solvation and Charge-Transfer Interactions Thermodynamic Characteristics The donor capacity of the halide ions is manifest with respect, not only to cations, but also to appropriate neutral molecules. With halogen or interhalogen molecules the products of this interaction are the chemically recognizable polyhalide ions (see Section 4), but with other, polar molecules, e.g. water, alcohols, amines or acetonitrile, the interaction typically involves a combination of ion-dipole, charge-transfer and dispersion forces, the balance of which governs the power of the liquid to solvate the halide ion. Most of our knowledge of the quantitative aspects of ion-solvation is restricted to water as the solvent. For the alkali-metal halides, for example, the measured free energy, enthalpy and entropy of the change M+(g)+X~(g) -> M + (aq) + X~(aq) indicate that in dilute solutions the thermodynamic properties of one ion are independent of those of an­ other293»295 ~297. This implies that the thermodynamic properties defining hydration can be split up into contributions from the individual ions. All of the numerous attempts which have been made to achieve this end involve some assumption that cannot be thoroughly substantiated, though most yield results which are in fairly close agreement. The absolute values given in Table 19 are based on estimates due to Halliwell and Nyburg295 and to Harvey and Porter 296 ; all of the values are relative to AG£yd = —252-0 kcal g-ion -1 , Ai/hyd = -260-7 kcal g-ion -i and S° = - 2 9 cal g-ion -i deg" 1 for the proton at 298°K. 294 L . Pauling, / . Chem. Soc. (1948) 1461. 295 H . F. Halliwell and S. C. Nyburg, Trans. Faraday Soc. 59 (1963) 1126. 296 K . B. Harvey and G. B. Porter, Introduction to Physical Inorganic Chemistry, Addison-Wesley, Reading, Mass. (1963). 297 F . D. Rossini, D. D. Wagman, W. H. Evans, S. Levine and I. Jaffe, Selected Values of Chemical Thermodynamic Properties, Circular 500, National Bureau of Standards, Washington (1952); National Bureau of Standards Technical Notes 270-1, 270-2, 270-3 and 270-4, U.S. Government Printing Office, Washington (1965-1969); Codata Bulletin, International Council of Scientific Unions, Committee on Data for Science and Technology (1970).

1236

CHLORINE, BROMINE, IODINE AND ASTATINE: A. J. DOWNS AND C. J. ADAMS

The absolute values of the therraodynamic changes accompanying the hydration of the Cl", Br~ and I~ ions appear to be primarily a function of ionic size. Thus AG£yd, A//£yd and S° all decrease numerically as the ionic radius increases, corresponding to the attenua­ tion of ion-dipolar forces and the diminishing restriction of the motion of water molecules round the ion. Following a theoretical interpretation first suggested by Born293, the free energy of electrostatic solvation of an ion of radius r' and charge z in a solvent of dielectric constant e is given by AGgoiv =

2r'

H

+1-32 kcal mol-i at 298°K

(3)

Of the several shortcomings of this treatment, it may be noted that the radius of an isolated ion in the gas phase or in solution is unlikely to be the same as that if) in a particular crystal lattice, while the macroscopic value of € is almost certainly inappropriate in the high electricfieldswhich operate close to the ions. If € is equated with the bulk dielectric constant of water, the results of Table 19 imply that for the halide ions r' exceeds r by about 0-2 Ä. On the basis of equations (2) and (3) Johnson293 has shown that the free energy of aqueous dissolution of an ionic halide MX» may be represented approximately by the relationship AG,°(kcal) =

256(/i+l)/i r++r-

164«2 + /iAG2 y d P"]-7-4«-6-l r++0-72

(4)

Expressed as a function of r +, AGS° passes through a maximum when r_ = r + [l-25V(l + l / " ) - l ] + 0 - 9 ( V ( l + l/H)

(5)

Being very approximate, equations (4) and (5) do not predict accurately the value of the cation radius at which the maximum in AGS° is achieved in a series of salts of given formula 4-20

4-10

ME MF

kcal mol"

-10

-20

-30

Csn 1 Sr2-t- Rb+ Cation radius, A

2

FIG. 19. Variations in the standard free energies of solution, AG£, of some alkali and alkalineearth metal halides with the crystal radius of the cation.

PROPERTIES OF THE HALIDE IONS

1237

type containing the same halide anion. Nevertheless, the following implications with regard to qualitative variations in AGS° are vindicated by experimental data (see Fig. 19): (i) For salts of the same anion, as r+ is increased, AGS° rises, reaches a maximum and then falls steadily towards a limiting value of («AG£ yd [X-]-7-4/z--6-l). (ii) For different series each con­ taining a fixed anion, AGS° reaches a maximum at a larger cation radius as the size of the anion advances. The maximum should occur when the sizes of the cation and anion are suitably matched, (iii) When two series of compounds have different formula types MX n but contain the same anion, the free energy maximum should occur at a larger cation radius in the series with the higher value of«. Provided that the phase deposited from the saturated solution is the anhydrous compound rather than a hydrate, these generalizations provide a useful and instructive guide to the solubilities of ionic halides, but do not apply well to the halides of metals like silver and mercury whose properties deviate markedly from those expected of ionic compounds. Optical Spectra There exist close analogies between the absorption spectra of ionic halides and those of the isoelectronic noble gases298. Thus, the first absorption band is attributable to transi­ tions of the type np6 -> np5(n + l)sl; there is some discussion about whether a weak absorp­ tion at slightly higher energy is due to a direct ionization of an «p-electron to the conduction continuum, or whether it corresponds to the Laporte-forbidden transition np6 ->np5(n+l)pl; further, strong absorption bands at somewhat higher wavenumbers can be ascribed to the transition np6 -> np$ndl. There is little doubt that the description in terms of transitions between localized atomic states is far more appropriate for the alkalimetal halide crystals than is an energy-band description. Because of the relatively localized electronic structure of many simple and complex metal halides it is possible to identify electronic transitions involving charge-transfer from the halide to the metal ion; recent years have witnessed extensive investigations of such transitions298. Although ion-dipole interactions undoubtedly play a major role in the solvation of halide ions, some orbital overlap and charge-delocalization are clearly implied by the optical spectra of the ions in solution. The ultraviolet absorption spectra of such solutions show intense (c ~ 104) bands, broad and structureless, which are considered to arise from a charge-transfer process: (X-)solv-^(X+e)solv

and which are thought to represent charge-transfer-to-solvent (c.t.t.s.) transitions299»300. Details of the ultraviolet spectra of aqueous solutions of the halide ions are presented in Table 19. Applying Mulliken's theory for donor-acceptor complexes (see pp. 1209-14) leads to the transition energy, £ m a x = Ix- - £ s o l v + Γ~ΖΈ~

(1/2r<)

(6)

depending on 7χ-, the ionization potential of the anion, 2ssoiv, the electron affinity of a group of solvent molecules, σ, an overlap term determined by the overlap integral and polari­ zation terms, and n9 the crystallographic radius of the ion. Good agreement is observed 298 Chr. K. Jorgensen, Halogen Chemistry (ed. V. Gutmann), Vol. 1, p. 265. Academic Press (1967); Progress in Inorganic Chemistry, 12 (1970) 101. 299 M . J. Blandamer and M. F. Fox, Chem. Rev. 70 (1970) 59. 300 M . F. Fox, Quart. Rev. Chem. Soc. 24 (1970) 565.

1238

CHLORINE, BROMINE, IODINE AND ASTATINE: A. J. DOWNS AND C. J. ADAMS

between predicted and measured transition energy values for halide ions both in acetonitrile and in aqueous solutions. An alternative approach relates EmAX to the heats of solvation ΔΖ730ΐν[Χ-] and Δ7/301ν[Χ] by Em&x = / x - + A # 3 0 l v [ X - ] - A / W X ] + X-i?

(7)

X representing the Franck-Condon strain arising from the immobilization of the solvent dipoles during the electronic transition and B the excited-state electron binding energy. The difficulty of calculating B has been overcome by Franck and Platzman, who assumed a spherically symmetrical excited state centred on the parent anion site: B is then the sum of two contributions, one arising from the polarization of the oriented solvent molecules around the ion, and the other from the electronic polarization of the solvent by the excitedstate electron. With these assumptions, equation (7) gives a good account of the observed Em&x values for the halide ions in aqueous solution but does not reproduce the effect on Em&x of such variables as change of solvent, added electrolyte, pressure and temperature. Other approaches are299»300: (i) a so-called "confined" model of an electron in a potential well, variations in the radius of which determine the changes in 2smax, and (ii) a "diffuse" model in which the variable radius parameter rd is combined with the Franck-Platzman treatment to predict both absolute values of 2smax and environmental effects. The consensus of the various theories is that the excited state characterized by the ultra­ violet absorption bands is either completely or predominantly defined by the solvent and centred on the parent anion site; it typically lies at an energy ca. 50 kcal or 18,000 c m - 1 above that required for ionization of X ~. Charge-transfer is reported to occur even with a supposedly "inert" solvent like carbon tetrachloride301. Measurements of optical spectra imply that the halide ions form weak complexes with molecular oxygen302 but relatively stronger complexes with sulphur dioxide as the acceptor303. The action of sulphur dioxide is presumably responsible for the moderate solubility of ionic halides in liquid sulphur dioxide. The excited state of the optical transitions also controls the photolysis of solutions of halide ions, which leads to transient radical-ions such as X 2 ~ and ultimately to products such as hydrogen and the free halogen300. The radicals formed in the primary photolytic process either recombine, dimerise or are scavenged by other solutes, e.g. N 2 0 , acetone or Brönsted acids. Charge-transfer spectra are also observed for crystalline and gaseous ionic halides304. The ultraviolet absorption spectra of crystalline or gaseous alkali-metal halides consist, for example, of one or more bands attributable to a transition which may be represented roughly as M + X"->MX

301 M. J. Blandamer, T. E. Gough and M. C. R. Symons, Trans. Faraday Soc. 62 (1966) 301. 302 H. Levanon and G. Navon, / . Phys. Chem. 73 (1969) 1861. 303 E. J. Wooc'house and T. H. Norris, Inorg. Chem. 10 (1971) 614. 304 L. E. Orgel, Quart. Rev. Chem. Soc. 8 (1954) 422. 305 Chr. K. Jorgensen, Orbitals in Atoms and Molecules, Academic Press, London (1962); Absorption Spectra and Chemical Bonding in Complexes, Pergamon, Oxford (1962). 306
PROPERTIES OF THE HALIDE IONS

1239

In terms of the simplest form of the Mulliken theory, the transition energy is given by the equation -Emax = ΙΧ~~ΕΜ++ΔΕ

= £χ-/Μ+Δ£

\

/

W

where / signifies the ionization potential and E the electron affinity of the species and ΔΕ is the change of interaction energy due to the charge-transfer. In keeping with this, the frequency of maximum absorption of the halides of a given metal follows the order of in­ creasing electron affinity, viz. iodine < bromine < chlorine. The general features are reproduced by many crystalline halides, the charge-transfer mechanism providing an effec­ tive interpretation of the colours of those halides formed by metals like copper, silver, lead and mercury, which are characterized by relatively high ionization potentials. Crystalline alkali-metal chlorides generally exhibit a single absorption maximum in the ultraviolet, the bromides a pair of bands separated by 3900-4800 c m - 1 and the iodides a pair of bands separated by 7700-9600 cm - 1 . The close correspondence between these doublet separations and the energy differences between the two lowest states 2P^/2 and 2P\/2 of the free halogen atoms (Cl, 882; Br, 3685; I, 7603 cm - 1 ) provides convincing support for the proposed mechanism, as does the fluorescence emission due to excited states of metal atoms, which has been observed for some gaseous halides305. As with the photolysis of halide ions in solution, chemical effects may also arise from charge-transfer absorption in a crystalline halide through secondary transformations which may compete with physical processes of deactivation of the excited state. Thus, the forma­ tion of metal atoms is responsible for the colour-centres produced by ultraviolet irradiation of alkali halide crystals and for the photosensitivity of the silver halides. Variations in the quantum efficiency of the production of colour-centres in alkali halide crystals indicate that the primary absorption process is virtually independent of temperature, but that the forma­ tion of metal atoms depends on the vibrations of the crystal lattice. Whereas the primary process of charge-transfer is completely reversible, only if some specific interaction takes place within the short lifetime of the excited state, is a metal atom produced. As the ground-state of a system incorporating a halide ion and acceptor species acquires an increasing degree of charge-transfer, the properties of the system become more and more specific, and any purely ionic model becomes progressively less realistic. Behaviour of Halide Ions in Solution: Nmr and Other Studies 306,307 The introduction of ions into the lattice characteristic of liquid water gives rise to changes which may be defined by structural and kinetic criteria. In the former case a coordination number may be derived to describe the most probable number of water molecules in the immediate environment of the ion. Although this is a convenient parameter for many transition-metal and other multi-charged ions, for the halide ions, as for the alkali-metal cations, the exchange of water molecules between the first coordination shell and bulk water is so rapid that it is neither meaningful nor expedient to define a separate solvated species in which the central ion has a specific hydration number. In these circumstances it is more appropriate to employ the second criterion and define the ion-solvent interaction in terms of the lifetime of a water molecule in the coordination shell. In contrast with the F - ion, Cl ~, Br ~ and I - in aqueous solution all belong to the "structure-breaking" category, tending to break down the water structure in their immediate neighbourhood and increase the mobility of adjacent solvent molecules,. Such behaviour

1240

CHLORINE, BROMINE, IODINE AND ASTATINE: A. J. DOWNS AND C. J. ADAMS

is implied by considerations of the entropy of solution, viscosity and temperature-dependence of ionic mobilities in alkali halide solutions. Thus, viscosities of dilute aqueous electrolyte solutions, η, can be represented by the equation308 ηΙη0= l+Aci + Bc

(9)

where ^0 is the viscosity of the pure solvent, c is the concentration and A and B are con­ stants. As a function of the interionic forces, A can be calculated theoretically, but the constant B, which takes into account ion-solvent interactions, is purely empirical and may have either a positive or negative sign. It being possible to separate B into individual ion TABLE 20. PROPEIUIES OF THE HALIDE IONS IN AQUEOUS SOLUTION*»1»

Property B coefficient in the Jones-Dole equa­ tion 17/170 = 1+Ac*+Bc (17 = vis­ cosity; c = salt concentration) (mol/l)-iat298°K Ratio of residence times of water molecules ΤΛ/Τ (ΤΛ = residence time adjacent to X" ion; τ = residence time in a site of the water lattice) at 294-7°K Parameter E in the Samoilov expres­ sion TA/T = exp(E/RT) (kcal/mol) at 294-7°K Ionic molal shift for H2O molecules, δ (ppm/mol) at 298°K 1 H resonance 17 0 resonance Ratio of the reorientational correlation times of H2O molecules τα1τα° (rd = correlation time for molecules close to X"; Td° = correlation time in pure water) at 298°K Activation energies for the reorienta­ tion of H2O molecules adjacent to halide ions, Er (kcal/mol) Nmr properties of halide nuclei: % Natural abundance Sensitivity at constant field relative t o i H = 1000 Ground state nuclear spin, h Expectation value p for the valence-shell /^-electrons (a.u.) Average excitation energy, Δ (Rydbergs)

-

00070

Br~

-

0042

1-

-00685

0-63

0-61

0-58

-0-27

-0-29

-0-32

-0075 -0-92

+0026 -1-35

+0053 -1-72

+0069 -2-35

2-3

0-9

0-6

0-3

4-3 19F 100 0-833 1/2 6-40 0-63 10-2

Paramagnetic shift of X"(aq) relative to X"(g), tfaq(ppm) (calculated)

ci-

F-

2-6-3-3 35C1 75-53 4-70 XlO-3 3/2

2-1-2-9

37C1 79ßr 8ißr 24-47 50-54 49-46 2-71 7-86 9-85 xlO-3 xlO-2 xlO-2 3/2 3/2 3/2 5-74 10-24 0-56 0-48 10-2 21-3 -225 -430

2-3-2-7 1271 100 9-34x10-2 5/2 12-76 0-40 31-9 -600

α C. Deverell, Progress in NMR Spectroscopy (ed. J. W. Emsley, J. Feeney and L. H. Sutcliffe), Vol. 4, p. 235, Pergamon (1969). b C. Hall, Quart. Rev. Chem. Soc. 25 (1971) 87.

PROPERTIES OF THE HALIDE IONS

1241

contributions, structure-breaking ions like Cl~, Br~ and I~ are identified by negative B coefficients, whereas structure-forming ions like F~ are identified by positive coefficients (Table 20). According to Samoilov309, ion-hydration can be considered in terms of the effect of ions on the translational motion of adjacent water molecules: the mean lifetime (rh) of a water molecule in close proximity to an ion is related to that (r) of a water molecule in an equilibrium position in the water lattice by the expression τΛ/τ = Qxp(EIRT)

(10)

Since τ is about 1·7χ 10-^ sec (at 21-5°C), the values of 0-58-0-63 calculated for rjr for Cl~, Br~ and I - indicate an extremely high frequency of exchange of water molecules from sites adjacent to an ion; further, the parameter E is negative, indicating an increase in the mobility of water molecules in the neighbourhood of the ions (see Table 20). This phenom­ enon has been termed "negative hydration". Solutions of halide ions are particularly suited to investigation by nuclear magnetic resonance techniques306»307. The resonating nucleus, whether of the halide ion or of an atom in the solvent molecule, undergoes rapid exchange between all possible environments in the solution; accordingly, only one signal is observed at an average frequency determined by the magnetic shielding and lifetime of the nucleus in each of the many allowed sites. The rapidity and randomness of ion-ion and ion-solvent encounters average local magnetic and electric fields to very small values to give relatively narrow resonance lines even for the quadrupolar nuclei of the heavier halogen atoms. Hence it is possible to detect compara­ tively small differences in magnetic shielding as well as broadening of the resonance line by chemical exchange, hyperfine interaction or quadrupolar effects. l

H and llO Resonances of Solvent Molecules™ Halide ions dissolved in water produce quite small but well-established changes in the l H resonance and relatively larger changes in the 1 7 0 resonance. A linear or near-linear variation of chemical shift with salt concentration is generally observed at concentrations typically up to 2 or 3 molal. The results are normally expressed in molal shifts, that is, the initial slopes of plots of *H or 1 7 0 chemical shift against molality. The ionic contributions appear to be additive at lower concentrations, and accordingly the shifts have been divided into individual ionic contributions. Proton chemical shifts produced by the halide ions other than F ~ arise primarily from the rupture of hydrogen bonds, in keeping with their structure-breaking propensities which follow the sequence Cl~ < Br~ < I - 3 1 ° . By contrast, direct ion-solvent interactions, rather than disruption of the solvent structure, appear to play the dominant role in determining 1 7 0 molal shifts. Whereas all the alkali cations produce similar effects on the 1 7 0 resonance of water, the halide ions show signi­ ficant differences amongst themselves in the effects they have on the 1 7 0 chemical shift. A very similar distinction between the behaviour of cations and anions in alkali halide solutions has been found in studies of their Raman spectra, which show the vibrational properties of the solvent molecules to be more strongly perturbed by the halide ions311. This result is somewhat surprising, especially since X-ray diffraction data indicate that, for the alkali-metal ions, the ion-dipole interactions with solvent water molecules occur via the oxygen atom, whereas it is commonly assumed (from energetic considerations) that the 3io J. C. Hindman, / . Chem. Phys. 36 (1962) 1000. 3Π G. E. Walrafen, / . Chem. Phys. 36 (1962) 1035.

1242

CHLORINE, BROMINE, IODINE AND ASTATINE: A. J. DOWNS AND C. J. ADAMS

halide ions interact via the centres of positive charge at the hydrogen atoms. Some Raman studies312 have been published in support of such a model for halide ion-water interactions. Possible explanations of the unusual response of the solvent molecules to halide ions invoke either a redistribution of charge in the H - 0 bonds induced by the polarizable halide ions or, alternatively, short-range repulsive forces which operate in collisions between ions and solvent molecules. Perhaps the most direct evidence of the "negative hydration" of the halide ions in aqueous solution has come from measurements of proton spin-lattice relaxation times306. In keeping with the enhanced mobility of water molecules close to Cl ~, Br - or I - ions, the proton relaxation rate is reduced as compared with that found in pure water. The variations follow closely the order expected from other estimates of the structure-breaking abilities of the ions; thus, the initial slopes of plots of relaxation rate against molality show an extremely close correlation with the entropy of solution for alkali halides. Halide Ion Resonances*06,»307

The nuclei 35C1, 79Br, 81Br and 127I give weaker resonance signals than protons with the result that most investigations are presently confined to solutions of concentration greater than 0-05 molal. On the other hand, a wide range of chemical shifts, spanning as much as 100 ppm, is observed as a result primarily of changes in the paramagnetic contribution to the magnetic shielding. Relaxation times and hence resonance linewidths of the heavier halogen nuclei are determined by processes originating mainly from the quadrupole moments of the nuclei. In consequence, investigations of linewidths can illuminate both the magnitude and time-dependence of electric-field gradients at the nucleus; being deter­ mined by the ion-solvent and ion-ion interactions in solution, such gradients give access to information about the times of rotational and translational motions of ions and solvent molecules. In aqueous solutions of alkali-metal chloride, bromide or iodide, the chemical shift (to low field) of the halogen nucleus varies with the concentration and with the identity of the counter-ion. For a given halide ion, the efficacy of the alkali-metal cations in producing low-field shifts may be ordered N a + < K+ < Li + < R b + < Cs + . It has been suggested that the ion shifts so induced originate in the direct binary interactions of the anion with the cation, the contribution from halide-solvent and halide-halide interactions being assumed approximately constant at all but the highest concentrations. Since the larger ions cause greater shifts than the smaller ions, the effect is consistent, not with a.direct electrostatic (polarization) interaction, but with the repulsive overlap of the closed-shell orbitals of the ions, a mechanism first suggested by Kondo and Yamashita to account for the shifts of the ions in alkali halide crystals313. The experimental shift may be expressed in terms of the fine-structure constant a, the expectation value „ of r{ ~3 for an outer/7-electron, and an average excitation energy Δ as -16**

r
δ = —_<ΓΓ3>2>[Σ Λ,.^+^.Η,Ο^-Λ,.^Ο

i 0

)!

(11)

where A w and A f _ Hj0 denote appropriate sums of the squares of overlap integrals for ion-ion and ion-solvent encounters. Table 20 lists for the halide ions values of p computed from Hartree-Fock wavefunctions and estimates of Δ derived from the ultra­ violet absorption spectra of solid alkali halides. Variations in the quotient Ο , - ^ ρ / Δ 312 G. E. Walrafen, / . Chem. Phys. 44 (1966) 1546. 313 J. K o n d o and J. Yamashita, Phys. Chem. Solids, 10 (1959) 245.

PROPERTIES OF THE HALIDE IONS

1243

account for a considerable part of the observed increase in the magnitude of the chemical shifts found with increasing atomic number of the resonating halide ion. Numerous measurements of the nuclear magnetic relaxation times of the halide ions in solution have been carried out. Though the theory of such relaxation314 allows us to relate the observed relaxation time to a function of a correlation time (a period of time giving a measure of the timescale of molecular motion), it is a major problem to determine the motion in the liquid to which the correlation time refers. The situation is further com­ plicated by the absence of any knowledge of the quadrupole coupling constant; more or less irrespective of the mechanism, this must itself be an average over many configurations, since the ion experiences many field gradients from various neighbours, all of them random func­ tions of time, notwithstanding at least some degree of correlation. It has been suggested that the solvent molecules contributing to relaxation are mainly those in the rather dis­ organized "outer-solvation" regions, characterized by structural mismatching between the bulk solvent and the solvent molecules oriented about the ions and by relatively free rotation or reorientation of solvent molecules. In certain cases the relaxation time of a quadrupolar halogen nucleus can be related to the rates of relatively fast reactions occurring in electrolyte solutions. 127 I, which typically exhibits the shortest relaxation time, can thus furnish kinetic information about processes with lifetimes of the order of nanoseconds. Thus, for the aqueous system I- +I 2 - l a ­ the broadening of the 127I resonance gives rate constants which support the view that the equilibrium is established at a diffusionally limited rate306»307; that the analogous reaction between Cl" and Cl2 exhibits a rate constant at 25 °C of 8 ± 4 x 106 sec "Us the inference drawn from a more comprehensive evaluation of the 35C1 and 37C1 resonance linewidths315. Like­ wise, analyses of the broadening of the 127 I~ and 81Br~ resonances in the presence of small concentrations of Hg 2 + , Zn 2+ or Cd 2+ ions have led to rate constants for processes such as Hgl 4 2-+i*- ^ i - + i * H g l 3 2 -

which also appear to be diffusion-limited. Having shown that Hg2 + ions bind to reactive or exposed sulphydryl groups of proteins, Stengle and Baldeschwieler316 have developed the so-called "halide ion probe" technique, wherein the mercury tag provides a possible coordination site for halide ions, with the result, in the event of rapid exchange, that the halide ion resonance is broadened. In principle this provides a means of determining the relaxation time of the halide ion nucleus at the binding site, and variation of the conditions, such as pH and temperature or the addition of other substances to the system, reveals changes in the accessibility of the binding site to the halide ions. Furthermore, the broad­ ening of the nmr signal may be used simply to monitor a titration experiment: for example, as mercuric chloride is added to a chloride solution containing small amounts of haemo­ globin, the variation of the 35C1 linewidth indicates two reactive -SH groups per haemo­ globin to which the mercury binds. Several other studies of this sort have been performed, and the technique has recently been extended to systems other than mercury-labelled mole­ cules, as in the investigation of the binding under various conditions of the Cl~ ion to a zinc metallo-enzyme carbonic anhydrase317. A survey of such studies of biomolecules has been given by Hall3<>7 314 H . 315 C . 316 T . 317 R .

G . H e r t z , Z. Elektrochem. 6 5 (1961) 2 0 ; K . A . Valiev, / . Exp. Theor. Phys. 10 (1960) 77. H a l l , D . W . K y d o n , R . E . R i c h a r d s a n d R . R . S h a r p , Proc. Roy, Soc. A318 (1970) 119. R . Stengle a n d J. D . Baldeschwieler, Proc. Nat. Acad. Sei. U.S.A. 55 (1966) 1020. L . W a r d , Biochemistry, 8 (1969) 1879.

1244

CHLORINE, BROMINE, IODINE AND ASTATINE: A. J. DOWNS AND C. J. ADAMS

3.2. GENERAL PROPERTIES OF HALIDES Attention is now directed to the following general aspects of compounds containing chlorine, bromine or iodine in combination with more electropositive partners: (i) classifi­ cation; (ii) methods of formation; (iii) physical characteristics, with particular reference to thermodynamic and spectroscopic properties; (iv) nature of bonding. Apart from the hydro­ gen halides, which form the subject of subsection 3.3, this is not the place to enter into detail about individual halides, the physical and chemical properties of which are presented under the heading of the appropriate electropositive component. It is with the halides as a particularly populous and important class of inorganic system that the present survey is concerned. The approach is dominated by considerations of the ligand function, real or formal, of the halide ions, and of variations in the nature of this function. Classification289,293,3i8-32i Concerning halides as a class certain generalizations may be remarked at the outset. First, primarily as a result of the small size of the F ~ ion and the strong bonds formed by fluorine with other elements, fluorides often differ in stoichiometry or in structure from the other halides. Thus, we note that many metal fluorides have three-dimensional lattices, whereas the corresponding chlorides, bromides and iodides form crystals having layer or

(c) FIG. 20. Representative structures of crystalline halides: (a) the NaCl and CsCl structures; (b) the Cdh layer structure [reproduced with permission from C. S. G. Phillips and R. J. P. Williams, Inorganic Chemistry, Vol. 1, p. 183, Clarendon Press, Oxford (1965)]; (c) unit cell of the structure of S11I4 [reproduced with permission from A. F. Wells, Structural Inorganic Chemistry, 3rd edn., p. 350, Clarendon Press, Oxford (1962)]. 318 N . V. Sidgwick, The Chemical Elements and their Compounds, Vol. II, Clarendon Press, Oxford (1950). 319 A . F. Wells, Structural Inorganic Chemistry, 3rd edn., Clarendon Press, Oxford (1962). 320 c . S. G. Phillips and R. J. P. Williams, Inorganic Chemistry, Vol. 1, Clarendon Press, Oxford (1965) 321 F . A . Cotton and G. Wilkinson, Advanced Inorganic Chemistry, 2nd edn., Interscience (1966).

GENERAL PROPERTIES OF HALIDES

1245

chain structures. Secondly, whereas many crystallinefluoridesand oxides are isostructural, chlorides, bromides and iodides are often structurally analogous to sulphides, selenides and tellurides of similar formula type, e.g. Pbl 2 and SnS2; Cdl2 and VTe2. Notwithstanding the substantial differences between the individual halogens, chlorides, bromides and iodides are sufficiently similar to permit a collective classification (see Table 21), though it must be appreciated that there are no clear lines of demarcation between the different classes. Rather is there a uniform gradation from halides which are for all practical purposes ionic, through those of intermediate character, to those which are essentially molecular. Structures representative of the different types of crystalline halide are illustrated in Fig. 20. Ionic Halides: Three-dimensional Lattices Ideally an ionic halide exists at normal temperatures as an involatile solid possessing a three-dimensional lattice in which the atoms are regularly disposed. Moreover, although the electrical conductivity of the solid is very low, it is markedly enhanced by fusion. Very careful examination by X-ray techniques has shown for several compounds of this type that the electron density distribution falls to a very low value between adjacent unlike nuclei, and that, if the distribution is divided at this point, each atom possesses an approximately integral net charge293. Such measurements disclose further that the internuclear distances are approximately additive. An ionic halide may also be defined as one whose lattice energy can be reproduced, within the limits of experimental error, by the Born-Mayer relation (equation (1)) or more elaborate versions of this expression. According to some or all of these criteria, the class of "ionic" chlorides, bromides and iodides may reasonably be taken to comprise halides of the type MX formed by alkali-metal and other large univalent cations, e.g. Cu+, Ag+, NH 4 + and NR 4 + , and of the types MX2 and MX3 formed by metals with large cations—alkaline-earth halides, PbCl2 and UC13. As with hydrides, the closest approach to ionic behaviour is found in the halides formed by elements at the extreme left of the Periodic Table. However, as compared with the hydrides, the pattern is altered because of the more favourable enthalpies of formation of the gaseous anions and because of the greater sizes and polarizabilities of the halide anions. Accordingly "ionic" halides have larger heats of formation and are thermodynamically more stable than the corresponding hydrides, while the elements which form them occupy a wider region of the Periodic Table. Semi-ionic Halides: Layer or Chain Structures and Systems Involving Metal-Metal Interaction The capacity of a cation M*+ to distort an ionic charge distribution depends on its polarizing power, which can be related to Z*/r2 (Z* = effective nuclear charge; r = ionic radius)322 or alternatively to the electron affinity of the ion 293 ; these properties reflect variations, not only in the ratio z/r, but also in the shielding abilities of the valence and core electrons [cf. the following values of Z*/r2 in a.u. (electron affinity in eV): K + , 0-26 (4-34); Cs + , 0-20(3-89); Cu+, 0-84 (7-72); Au+, 0-86 (9-22)]. As the polarizing power of the cation increases, non-coulombic interactions make progressively larger contributions to the bond­ ing of metal halides. In the sequence KC1, CaCl2, ScCl3, T1CI4, for example, we pass from a compound well represented by the ionic model to a molecular species most aptly described by some form of bond model. The size and polarizability of the halide ion are also important in determining the character of the halide. Thus, whereas the majority of metal difluorides, which adopt the rutile or the fluorite structure, give a good approximation to the ionic 322

R. B. Heslop and P. L. Robinson, Inorganic Chemistry, 3rd edn., Elsevier (1967).

MIn, order dictated by lattice energy.

Generally MFn > MCln > ΜΒΓΛ >

High« Generally MF„ > MCln > MBrn > MI«, order dictated by coulombic interactions. High«

Boiling point

Melting point

Comparable with binding energy of atomic units. Commonly vapor­ ize to give polymer units.

Well represented by the ionic model: regular, three-dimensional lattices involving high coordination num­ bers; internuclear distances addi­ tive; lattice energies well repro­ duced by the Born-Lande or Born-Mayer equations.

Majority of pre-transition metals, lanthanides and actinides in lower charge states (+2, +3)

Ionic halides

Sublimation energy

Electronegativity difference between MandX Description

Formed by

Property

Non-metals and metals in high oxi­ dation states ( > + 3)

Molecular halides

Intermediate

Intermediate

May vaporize to give polymer units.

Increasing

Low Typically MIn >MBrn >MCl n > MFn, order dictated by polarizability. Low Typically MIn >MBr« > MCln > MFn, order dictated by polarizability.

Much smaller than binding energy of atomic units. Molecular units common to the solid and vapour phases.

Most conveniently described in terms Best represented by the bond model of departures from the idealized involving, for example, valenceionic model; some systems amen­ bond or molecular-orbital accounts. able to treatment by the band Stereochemistry of molecules can be treated in terms of valence-shell model. Halogen commonly found in relatively unsymmetrical en­ electron-pair repulsions. Weak vironment and lattices lack the intermolecular interactions due to regularity of the idealized systems, London, dipole-dipole, dipolee.g. layer or chain structures. Inter­ induced dipole and quadrupolar nuclear distances less nearly addi­ forces. tive. Non-coulombic interactions contribute significantly to the lattice energy.

Increasing

i/-Block transition metals in low charge states (+1 to +3) and Bmetals

Macromolecular or 'semi-ionic' halides

TABLE 21. PHYSICAL CHARACTERISTICS OF HALIDES MX n 318 " 322

Hydrolysis

Solubility

£ Heat of formation, — ΔΑ/,2980κ per ^ g atom of halogen

Solid

AH/ poorly reproduced by calcula­ Ionic model qualitatively and quan­ tions based on the simple ionic titatively inappropriate. Variations model. Deviations increase in the of AHf determined principally by order MFn but not reversed, by such deviations. M-Cl > M-Br > M-I.

Very low conductivity.

Very low conductivity presumably due to auto-ionization.

Inoreasing tendency

Favoured by polar, coordinating solvents of relatively high dielectric con­ Solubility determined by weak van der stant. Dictated by balance between lattice energy and solvation energies Waals' and dipolar forces between of the ions (see p. 1236). molecules. Dissolution commonly + For a given cation, e.g. K , solu­ Non-coulombic forces tend to stabi­ favoured by less polar media, e.g. bility in water typically increases lize the solid lattice with respect to benzene or CCI4. in the order MFn
AH/ well reproduced by calculations based on the simple ionic model. Variations of AH/ determined principally by variations of lattice energy.

Conduction may be facilitated (i) by ion-diffusion through relatively disordered lattice, e.g. α-AgI, or (ii) by electronic transport, e.g. Hgl2.

Low conductivity because of high activation energy for ion-diffusion.

General trend towards hierher values

Relatively low conductivity.

High conductivity.

Transitions observed as broad bands Transitions observed as discrete, of lower energy; in some cases there rather narrow absorption bands of may be an absorption edge beyond relatively low energy. which absorption is continuous. Transitions probably delocalized over the whole crystal. Photoconduction may occur, as in Hgl2.

Transitions observed as discrete, rather narrow absorption bands of high energy and probably highly localized.

Charge-transfer properties

Electrical conductivity Melt

Colour, if any, due to charge-transfer Colour, if any, usually due to chargeprocesses of the type M+X~!%! transfer transitions. MX· and/or to internal transi­ tions of cation.

Colour, if any, characteristic of in­ ternal transitions of cation.

Colour

1248

CHLORINE, BROMINE, IODINE AND ASTATINE: A. J. DOWNS AND C. J. ADAMS

model, other metal dihalides are mostly characterized by layer structures of the CdCl2 or Cdl 2 type, and the discrepancies, approaching 20% in some instances, found between the lattice energies calculated from equations such as (1) and those derived from a Born-Haber cycle allude to considerable non-coulombic interactions323. Simultaneously it is found that internuclear distances depart significantly from the additivity rule. These findings are attributable to the peculiarities of the layer structures in which the anion has only three nearest cation neighbours, all of them on the same side of it, an unsymmetrical arrange­ ment clearly incompatible with the requirements of the simple ionic model. The layers are held together by van der Waals' forces between adjacent halide ions. Reliable data con­ cerning the lattice energies of trihalides are sparse, but it is again noteworthy that, in contrast with trifluorides, which favour relatively regular structures incorporating octahedral MF<5 groups, other metal trihalides typically involve the layer structures characteristic of CrCl3 or Bil 3 319. Irregular environments are also encountered for metal ions like Cu2 + which lack cubic symmetry in their electronic ground states, and for B-metal ions like Tl + which possess low-lying excited states of less than cubic symmetry324. Whereas hydrogen gives rise to an extensive intermediate class of metal-like or inter­ stitial hydrides, the size and orbital energies of the halogen atoms largely preclude their entry into alloy-like structures. The nearest approach to such structures is found in the iodides of low-valent metals, such as Til 2 and Nbl 4 , which, with gross defect structures, behave as semiconductors or metals320. Semiconductor behaviour is also found for heavymetal halides with more regular structures, the conductivity depending either on ionmigration (e.g. Agl) or on electronic transport (e.g. Hgl 2 ). Such halides are amenable to descriptions based, at least qualitatively, on the band model; typical band gaps (in eV) are: CuBr, 2-9 and Agl, 2-8. Further, there are now known many halide species which incorpor­ ate clusters of metal atoms held together by strong metal-metal bonding, e.g. [Re 2 Cl 8 ] 2_ , [Re 3 Cl 12 ] 3 -, ReCl3, ReCl4, [M 6 X 12 ] 2 + (M = Nb or Ta; X = Cl, Br or I) and [M'6X8]4 + (Μ' = Mo or W; X = Cl, Br or 1)325-327. The formation of these systems, characterized by interatomic distances close to those found in the parent metal, appears to be favoured by metal atoms having relatively unfavourable ionization potentials for the removal of 1-3 electrons, usually coupled with a large heat of atomization. In fact, most binary metal halides other than those of transition metals in oxidation states > + 3 are most profitably discussed in terms of the simple ionic model and of devia­ tions from this model. Features of the halides which can thus be rationalized include (a) the thermodynamic functions defining the formation of the compound289»293, (b) the suscepti­ bility to redox reactions of the type MX„ + £X2 ^ MX W+1 (see pp. 1119-20P 9 ' 293 , (c)the capacity to undergo halogen-exchange reactions of the type R - Y + M X ( s ) - > R - X + MY(s)289'293, (d) the stability of halides containing complex cations susceptible to thermal decomposition, e.g. PH 4 + X ~, which commonly decreases in the order I > Br > Cl > F 293 , and (e) the solubility in various media (see pp. 1236-7)293. In each case the behaviour of the halide is principally a function of the lattice energy, major variations of which reflect variations of ionic charge and size. There is ample evidence that the simple ionic model is 323 D . F . C . Morris, / . Inorg. Nuclear Chem. 4 (1957) 8. 324 j . D . Dunitz and L. E . Orgel, Adv. Inorg. Chem. Radiochem. 2 (1960) 1. 325 F . A . Cotton, Quart. Rev. Chem. Soc. 2 0 (1966) 389; Rev. Pure Appl. Chem. 17 (1967) 25. 326 D . L. Kepert and K. Vrieze, Halogen Chemistry (ed. V. Gutmann), Vol. 3, p. 1. Academic Press (1967); M . C. Baird, Progress in Inorganic Chemistry, 9 (1968) 1. 327 j . H . Canterford and R. Colton, Halides of the Second and Third Row Transition Metals; WileyInterscience (1968).

GENERAL PROPERTIES OF HALIDES

1249

in at least qualitative accord with the energy trends of numerous chemical processes even when there are marked deviations from the properties normally associated with "ionic" behaviour. As pointed out elsewhere293»320j this is a tribute, not to the descriptive accuracy of the model, but to the self-compensating features embodied in setting up an ionic charge distribution. As the size of the halide anion changes, the response of the lattice energy varies markedly from one series of halides to another. Thus, with increasing anion size, the lattice energies of the monohalides of copper, silver and gold diminish much less rapidly than do those of the alkali metals, while those of the dihalides of cadmium, mercury and lead fall more slowly than do those of alkaline-earth metals of comparable size293. This implies that, relative to either the alkali-metal or alkaline-earth series, there is an increase in the strength of the cation-anion binding as the atomic number of the halogen increases. The effect is clearly + 400

6 +300

+200

Ϊ I i

+100l·

FIG. 21. Differences between the heats of formation of A-metal and B-metal halides for cations of similar size and the same formal charge [after C. S. G. Phillips and R. J. P. Williams, Inorganic Chemistry, Vol. II, p. 503, Clarendon Press, Oxford (1966)].

illustrated in Fig. 21 by comparing the heats of formation of A-metal and B-metal halides, MAXw and MBX» respectively, for cations of the same formal charge and of similar size320. According to the simple ionic model, the differences between the heats of formation of two

1250

CHLORINE, BROMINE, IODINE AND ASTATINE: A. J. DOWNS AND C. J. ADAMS

such gaseous cations [Δ#/(Μ Β Λ+ )-Δ77/(Μ Α Λ+ )] should be reproduced in the differ­ ences between the heats of formation of analogous halides [ Δ # / ( Μ Β Χ ? Ι ) - - Δ 7 7 / ( Μ Α Χ » ) ] · That the latter quantity is invariably more exothermic than the former, the divergence increasing in the sequence F - < Cl ~ < Br ~ < I ~, betokens the effect of polarization or partial covalent bonding between the B-metal and halide ion. As a result, B-metal ions tend to be class b acceptors328, forming complexes with stability constants in the order I - > B r - > C l ~ > F _ rather than in the reverse sequence characteristic of the metals whose halides approximate better to the ionic model. Molecular Halides The halides of non-metals and of metals in high oxidation states usually take the form of discrete molecules which persist throughout the solid, liquid and gaseous phases, e.g. Snl4, SbCl5, WC16 and Al2Br6. The molecules are held together by van der Waals' attrac­ tions, primarily between the halogen atoms of different molecules, which may be augmented by dipole-dipole and dipole-induced-dipole forces. Accordingly, such halides are charac­ terized by relatively low lattice energies, by low melting and boiling points, by solubility in non-polar solvents and by very meagre electrical conductivities in both the solid and liquid states. For ionic halides, the boiling and melting points tend to fall as the atomic number of the halogen increases, the work required to separate the ions diminishing as they become larger: for molecular halides, the boiling and melting points tend to rise with the atomic number of the halogen (see, for example, the hydrogen halides, subsection 3.3), the polarizability of the halogen here exercising the dominant influence. It must be emphasized, however, that these trends represent the limiting behaviour of each class. Again, whereas most ionic halides dissolve in water to give hydrated metal and halide ions, molecular halides generally suffer ready and irreversible hydrolysis. Within a halide molecule of the type MXn, at no point between the M and X nuclei does the electron density approach zero. To account for the electron distribution, we must consider as a primary influence σ-bonding between M and X, and, as a secondary influence of somewhat uncertain importance ττ-bonding, viz. M ^= X (e.g. ρπ -> dn interaction, as in SiCl4) or M =^ X (e.g. dn -> d% interaction, as in PtCl62 _ ). The stereochemistry of the aggre­ gate is best interpreted on the basis of the bond model, whether embodied in localized molecular-orbital methods329 or in the principle of repulsion between the valence-shell electron pairs of the central atom330. For many practical purposes the average bond energy B(M — X) = Ι/ηΔΗ, where ΔΗ is the enthalpy change for the process MXw(g) -> M(g)+«X(g), is most easily related to the thermodynamic properties of MX n . If the assumption of additivity is a reasonable approxi­ mation, it is thus possible qualitatively to account for variations of thermal stability, for example, with respect to a reaction such as MXn(g) -> M X n - l ( g m X 2 ( g )

or for the occurrence of substitution reactions, e.g. MXn(gmH2(g) -> MXn-i(g)+HX(g) For a given value of«, B(M — X) invariably decreases in the order X = F > Cl > Br > I, which is also the sequence of stability of MXW with respect to the elements. In the 328 s . Ahrland, J. Chatt and N . R. Davies, Quart. Rev. Chem. Soc. 12 (1958) 265.

329 R. G. Pearson and R. J. Mawby, Halogen Chemistry (ed. V. Gutmann), Vol. 3, p. 55. Academic Press (1967).

330 R. J. Gillespie, Angew. Chem., Internat. Edn. 6 (1967) 819.

GENERAL PROPERTIES OF HALIDES

1251

interpretation of chemical changes, the bond energies of molecular halides fulfil a function analogous to that of the lattice energies of ionic halides. However, the bond-energy ap­ proach is marred by the fallibility of the additivity principle and by the fact that bond energies, unlike lattice energies, are not easily accounted for by the theoretical models presently available. Halide Complexes In common with fluoride, chloride, bromide and iodide ions function as ligands to form, with the majority of metal ions or molecular halides, complexes such as AgCl2 ~, FeCl4 -, SbBr4 -, Hgl 4 2 -, WC162 ~, etc., as well as a wealth of mixed complexes in company with other ligands; this second class includes neutral or charged species, e.g. [Co(NH3)4Cl2]+, NbOCl 4 -, [Cl5RuORuCl5]4-, CoBr2(PHPh2)3, [Cul(bipyridyl)2] + Mn(CO)5I and [^-C 2 H 4 )PtCl 3 ] -. In the solid state there exist the following possibilities for a compound of the formula A m BX» 319 : (a) A, B and X form an infinite three-dimensional array of ions in which no finite or infinite one- or two-dimensional complex can be distinguished, e.g. complex fluorides of the type ABF 3 ; (b) B and X form infinite two-dimensional or one-dimensional complexes, as in NH 4 CdCl 3 ; (c) B and X form finite complexes, as in K2SeBr6, (NH 4 ) 2 PdCl 4 or R 4 NHX 2 (R = Me, Et or Bu); in some cases finite polynuclear complex ions are en­ countered, e.g. T12C193-, [Mo 6 X 8 ] 4+ and [Nb 6 X 12 ] 2+ . Sometimes, too, a compound AmBXn may contain, not ions BXW, but ions BX^-p and/? separate X~ ions, as in the case of Cs3CoCl5, which should accordingly be written Cs + 3 [CoCl 4 ] 2 -Q-. Again, there may be complex ions of two types, as in PC15 containing the ions PC1 4 + and PC1 6 _ and (NH4)2SbBr6 containing SbBr63 - and SbBr6 - ions. The combination of a univalent metal halide AX and a molecular halide BXn-m to give a finite complex BXnm ~ can be represented in the form of a thermodynamic cycle: mAX(s)

+

BX n _ m (g)

mA+(g) + mX~(g) + BXn.m(g)

—

-

-A m BX n (s)

- mA+(g) + BX n m ~(g)

The stability of AmBXw then hinges on whether the enthalpy of interaction (x) of gaseous X _ with gaseous BXn -m to give gaseous BXnm ~ compensates for the decrease of lattice energy which attends the change from the simple to the complex halide. Since the differ­ ence in lattice energy decreases with advancing size of the A + cation, it follows that the condition of stability is most likely to be fulfilled when the cation is large, a conclu­ sion in agreement with the observation that caesium salts are the most stable of the alkali-metal salts with respect to thermal decomposition of anions such as HC1 2 ", BC14~, TiCl62~ and i 3 - 289,331. The same principle is evident in the extensive use made of quaternary ammonium, phosphonium or arsonium species or of large organic cations derived from amines to stabilize polyhalide, hydrogen dihalide and MX 4 2 - anions (M = first-row transition metal). Conversely the energy of formation and large size of the complex halogeno-anions can be turned to advantage in the preparation and stabilization of less common cations; examples include Cd 2 2 + as [Cd2][AlCl4]2 3*2, N O + and N 0 2 + as 331 F . Basolo, Coord. Chem. Rev. 3 (1968) 213. 332 j . D . Corbett, Inorg. Chem. 1 (1962) 700.

1252

CHLORINE, BROMINE, IODINE AND ASTATINE: A. J. DOWNS AND C. J. ADAMS

[NO]2[SnCl6] and [N0 2 ][SbCl 6 pw, ici 2 + as [IC12][A1C14] and [ICl2][SbCl6pi9, and AsCl4 + as [AsCl4][SbCl6]333. It is noteworthy that none of these entities forms a salt with simple halide ions, implying that the part played by the energy liberated in forming the complex anion must be a decisive one. In this respect, chloro-complexes appear to be decidedly inferior to fluoro-complexes such as BF 4 ~, SbF6 - and PtF 6 ~, which have proved particularly effective in the stabilization of non-metallic cations289. However, that the formulation of the solid complex in terms of discrete ions may often be a severe over­ simplification is all too clearly demonstrated by the crystal structure of ICl2SbCl6 (see p. 1351)319. A survey of the relative stabilities of complex halide ions with respect to dissociation into their constituent ions in aqueous solution distinguishes entities such as H+, Be2 + , Ce 3+ , Fe 3 + and Ti 4 + (class a), which form their most stable complexes with F~, from entities such as Pt2 +, Pt4 + , Ag + , Tl3 + and Hg2 + (class b), which form their least stable complexes with F~ and their most stable complexes with I - 289,328,334,335. No rigorous theoretical explanation for either sequence or for the existence of the two classes of acceptor relative to the halide ions has been given. It is likely that the polarizing power Z*/r2 (see p. 1245), polarizability and the ability of the metal and ligand to engage in M -+Χ(απ-άη) bonding (an opportunity denied to fluorine) are all significant factors, but their relative importance remains unfathomed. However, the stability of a complex ion in solution depends not only on the absolute strength of the M-X bond but also on the relative solvation energies of all the species involved. Hence, the balance of an equilibrium such as MF62 - + 6C1" ^ MCI62 - + 6F -

turns not only on the balance of M-F and M-Cl bond energies but also on the difference between the hydration energies of MF 6 2 - and MC162 ~ and six times the difference in hydration energies of F - and Cl ~. Since the hydration energies of F ~ and Cl ~ differ by no less than 33 kcal, the order of stability constants is not necessarily that of M-X bond energies in the complex ions. Indeed, the few available results suggest that even for complexes having stability constants which decrease in the sequence I > Br > Cl > F, the actual order of bond energies is F > Cl > Br > I. What seems to be the critical factor about the class b acceptors is that the M · · · X interaction energies, in common with the lattice energies and heats of formation of related binary halides, decrease less markedly as a func­ tion of the atomic number of X (see p. 1249). To this extent, the distinction between the two classes of acceptor is somewhat less fundamental than has been represented elsewhere289»32o,328,336. When it comes to isolating a complex from solution, solubility considerations play a major, sometimes a dominant, part; there are known many instances in which the species preponderant in solution is not the one to be precipitated from that solution. The presence of complex ions which are very much larger than the counter-ions means that lattice energies are relatively low and vary but little with the size of the counter-ion; by contrast, the major contribution to the solvation energy is made by the counter-ion, being roughly inversely proportional to the radius of that ion. Hence, with increasing size of the counter-ion, the 333 v . Gutmann, Z. anorg. Chem. 266 (1951) 3 3 1 ; F . J. Brinkmann, H . Gerding and K. Olie, Rec. Trav. Chim. 88 (1969) 1358. 334 L . G. Silten and A . E. Martell (eds.), Stability Constants of Metal-Ion Complexes, Chemical Society Special Publication N o . 17 (1964); Supplement N o . 1, Chemical Society Special Publication N o . 25 (1971). 335 G. P. Haight, jun., Halogen Chemistry (ed. V. Gutmann), Vol. 2 , p . 351. Academic Press (1967). 33ό A . J. Poe and M. S. Vaidya, / . Chem. Soc. (1961) 1023.

GENERAL PROPERTIES OF HALIDES

1253

situation arises that the sum of the solvation energies of the ions taken separately decreases more rapidly than the lattice energy of the solid complex. Awareness of this generalization is often useful in seeking to isolate a particular halide complex. Thus, of the alkali metals, caesium tends to form the least soluble salts of complex halide anions, a feature which has been exploited in the characterization of the RhCl62 - ion293. However, for complexes of very large cations, e.g. Et 4 N + and Bu 4 N + , a much larger proportion of the solvation energy is contributed by the anion, with the result that the solubility increases again (cf. Fig. 19). Accordingly, whereas M2[RhCl6] (M = Cs or NMe4) has been isolated from aqueous media, corresponding derivatives of the cations K +, Rb + , NEt 4 + and NBu4 + have eluded attempts to synthesize them293. Finally it may be mentioned that in effecting the separation of metal ions for analytical or radiochemical purposes, advantage has frequently been taken, in conjunction with ionexchange or solvent-extraction procedures, of equilibria involving the formation of halide complexes 337-341 . For example, the marked difference in stabilities of the chloride com­ plexes formed by Co 2 + and Ni 2 + in aqueous solution makes possible the efficient separation of these two ions by elution on an anion-exchange column with concentrated hydrochloric acid321. Likewise, the separation of actinides as a group from the lanthanide ions, together with partial fractionation of the actinides, can be effected by the use of 10 M lithium chloride solution to elute the ions from a suitable anion-exchange column operating at elevated temperatures (up to ca. 90°C)340»342. Again, the stable halide complexes characteristic of the B-metals are a party to the selective extraction of these metals from strongly acidic aqueous solutions of the halogen acids into organic solvents such as ethers, higher alcohols, ketones or t-butyl phosphate337»338»341. General Preparative Methods32**"322»327»343~345

The general methods of preparing binary chlorides, bromides and iodides, summarized in Table 22, may be divided into two classes, namely wet methods, which are feasible when the halide is not hydrolysed, and dry methods, which are obligatory when such hydrolysis occurs. A further distinction may be made between those methods where oxidizing condi­ tions prevail and those involving reduction of a higher halide, whether by external means or by spontaneous thermal decomposition or disproportionation. Of the oxidizing methods, those involving direct interaction of the elements and halogenation of an oxide are the most widely applicable. In the direct halogenation of an element, it is notable that the chloride produced may contain the element in a higher oxidation state than does the bro­ mide or iodide. However, it is commonly possible, within certain limits, to influence the nature of the product by choice of such reaction conditions as temperature or reactant 337 H . M . N . H . Irving, Quart. Rev. Chem. Soc. 5 (1951) 200. 338 D . F . Peppard, Ado. Inorg. Chem. Radiochem. 9 (1966) 1. 339 F . Helfferich, Ion Exchange, McGraw-Hill (1962). 340 j . Korkisch, Modern Methods for the Separation of Rarer Metal Ions, Pergamon (1969). 341 Y . Marcus and A . S. Kertes, Ion Exchange and Solvent Extraction of Metal Complexes, WileyInterscience (1969). 342 j . j . Katz and G. T. Seaborg, The Chemistry of the Actiniae Elements, Methuen, London (1957). 343 R. Colton and J. H . Canterford, Halides of the First Row Transition Metals, Wiley-Interscience (1969); D . Brown, Halides of the Lanthanides and Actinides, Wiley-Interscience (1968); K. W. Bagnall, Halogen Chemistry (ed. V. Gutmann), Vol. 3 , p . 303. Academic Press (1967). 344 j . D . Corbett, Preparative Inorganic Reactions (ed. W. L. Jolly), Vol. 3, p. 1. Interscience, N e w York (1966). 345 z . E. Jolles (ed.), Bromine and its Compounds, Benn, London (1966).

A

Halogenation of metal oxides

3

(ii) with molecular halogen in the presence of carbon (iii) with carbon tetrahalide or other organic halide

(i) with molecular halogen

Metal + hydrogen halide

Direct interaction of the elements

DRY METHODS

2

1

1

Method

Ti0 2 + C+2C12 Ta2Os+C+5Br2 Sc 2 0 3 + CCU Nb 2 0 5 + CBr 4 1

°2+Br2

°2Br2

> ττη,

W

t^nor > ™* £-—> 2TaBr5 370 ? C >> NScCl bBr 3 5 r Iflux

———>ZrCl4

>Pul3+3/2

400-600°c »

Pu+3HI 250-300 c

TJOJ I Π if" — ίΎΊΡΓΊ i

1

H2

2

Cr+2HC1 _ ^°° C > CrCl2 + H2 Co+2HBr2f*h™' > CoBr2+H2

300oC ΜθΙ 3

Zr02+2C12 W

MoCl5

Μ Ο ^ Μ Ο Β Γ , ^ * Μ Ο Β Γ

cu^

2Al+3Br 2 -*Al 2 Br 6 Sn+2C1 2 ->SnCl4 2Ta+5I 2 ->2TaI 5 Th+2C1 2 ->ThCI 4

Examples

Particularly useful and effective method of making metal chlorides.

Useful for chlorides and bromides.

Elevated temperatures invariably required. Oxyhalides may be formed. Useful for chlorides and bromides.

Useful method of preparing conventional lower-valent chlorides or bromides (cf. MoCl2 or Ta6Cli4).

Perhaps the most important general method. Elevated temperatures usually required, although with transition metals rapid reac­ tion can often occur with Cl2 or Br2 when THF or other ethers are used as the reaction medium, the halide being obtained as a solvate. Where different oxidation states are possible, chlorine, at elevated temperatures, tends to give a higher state than bromine or iodine. By choice of reaction conditions (e.g. temperature, reactant proportions, etc.) it may be possible to preselect the halide produced, e.g. PCI3 or PCI5.

Comments

TABLE 22. MORE IMPORTANT GENERAL METHODS OF PREPARING BINARY CHLORIDES, BROMIDES AND IODIDES»

to

Reduction of higher halide (i) with the parent element

6

1

Halogen exchange

. 400 600 c

-kPrPtri 1 "RfliK

°

410 C

►SAICI

■ > 3TaI5 + 5AlCl3

Ta

l 5 + T a «α-575-c ? T a 6 l i 4 Thermal-gradient technique particularly efFective.d-e

A1C13+2A1

3TaCl5 + 5AlI3

MCl3 + 3HBr - ° >MBr3 + 3HCl (M = lanthanide or Pu)

Ff^fln rlVRri

An excess of one reagent is usually required, equilibria normally being established.

A method which has been applied relatively widely. The acetone and/or methanol often becomes coordinated to the halide, but can usually be removed by gentle heating or pumping.to

CuCl2,2H 2 O HC1>150 ° C > CuCl 2 +2H 2 0

(iii) in presence of anhydrous hydrogen halide or mixed with ammonium halide (iv) with 2,2-dimethoxypropane MX„,mH 2 0+wMeC(OMe) 2 Me ->MX n + mMe 2 CO+2mMeOH

Efficient method of preparing certain anhydrous chlorides.

[Cr(H 2 0) 6 ]Cl 3 + 6SOCl2 reflux ) CrCl3 + 12HCl+6S0 2

(i) in vacuo

(ii) with thionyl chloride

Risk of oxyhalide formation.

Has been used to prepare trichlorides of the lanthanides.

This method works well for lanthanide and actinide tribromides.f

6SO2

> 8T n O i 1

> WOCU

AII3 has been used extensively for the pre­ paration of binary transition-metal iodides (the metal sulphide has also been used in place of the oxide). d,e Final product commonly an oxychloride.

[M(H 2 0) 6 JBr 3 7 ° 17 ° C > MBr3 + 6 H 2 0 Controlled vacuum thermal decomposition, M = lanthanide or actinide element.

Dehydration of hydrated halides

5

4

WO3+SOCI2 —reflux 4Lu 2 03+9Cl2+3S 2 Cl2

(v) with thionyl chloride Other halogenating agents include the hydrogen halide, NH4X, C12 + S2C12, COCI2 and PCI5

— 2 3 0 ° c ■ > Mol 2

M0O2+AII3

(iv) with aluminium trihalide

B

Hydrolysis of the molecular halogen

3

> MVu 1 H Y

> TaBr3 + TaBr5

°

500 c

AuCl3 2TaBr4

-► [Fe(H20)6]Cl2

3X2 + 6OH- ->5X"+X0 3 -+3H 2 0

Ag + +Cl-->AgCl Cu2++2I" ->CuI+iI 2

CoC03+2HI -> [Co(H20)6]I2

Fe+2HC1

>AuCl+Cl2

°

160 c

°

> MoI 2 +il2

i00 c

MoI3

ReCl 5 a t b p · > ReCl3 + Cl2

3WBr5 + Al475"240°C> 3WBr4+AlBr3 thermal gradient

(M = Sm, Eu or Yb; X = Cl, Br or I)

M Y , i 1H!

Examples

Used commercially for the production of alkali-metal halides.

Evaporation of the solution affords the halide, commonly in the form of a hydrated solid. Partial hydrolysis may give oxyhalides, e.g. BiOCl.

Feasible when the halide is not hydrolysed.

These are the principal methods of preparing binary halides of elements in lower valence states which are avoided under oxidizing conditions.

Comments

Gmelins Handbuch der Anorganischen Chemie; P. Pascal, Nouveau Tratte de Chimie Minerale, Masson et Cie, Paris. F. A. Cotton and G. Wilkinson, Advanced Inorganic Chemistry, 2nd edn., Interscience (1966). Z. E. Jolles (ed.), Bromine and its Compounds, Benn, London (1966). R. Colton and J. H. Canterford, Halides of the First Row Transition Metals, Wiley-Interscience (1969). J. H. Canterford and R. Colton, Halides of the Second and Third Row Transition Metals, Wiley-Interscience (1968). D. Brown, Halides of the Lanthanides and Actinides, Wiley-Interscience (1968). * P. M. Druce, M. F. Lappert and P. N. K. Riley, Chem. Comm. (1967) 486.

a b c d e f

Precipitation reactions

(ii) of the oxide, hydroxide or carbonate

WET METHODS Dissolution in aqueous halogen acid (i) of the metal

Thermal decomposition or disproportionation of a halide

(iii) with aluminium

Reduction of higher halide (cont.) (ii) with hydrogen

Method

2

1

7

6

Table 22 (cont.)

1257

GENERAL PROPERTIES OF HALIDES

proportions. Another useful route to certain anhydrous halides of metals in lower oxidation states involves the dehydration of the hydrates readily isolated from aqueous solution. Reduction or decomposition of higher halides provides a route to halides like InCl or Ta 6 I 14 which are avoided by reactions characterized by oxidizing conditions344; in practical terms the reaction is often conveniently effected by means of the thermal-gradient technique327, which is particularly suited to the preparation of halides having a narrow stability range. For mixed halide species, e.g. oxyhalides like W0 2 Br 2 346 or carbonyl halides like Mn(CO)5X347, the most widely applicable methods of preparation involve (a) substitution of a binary halide, e.g. with CO, C 5 H 5 -, CN -, OH - or 0 2 , or (b) halogenation, e.g. of an oxide, sulphide or carbonyl. Halide-exchange, so prominent as a method of producing fluorides, is less important to the preparation of the heavier halides, except where specific properties, e.g. volatility or solubility, facilitate exchange: sodium iodide in acetone, for example, is a good reagent for the replacement of chlorine by iodine mainly because sodium iodide is soluble and sodium chloride insoluble in this medium289. For the isolation of solids containing complex halide ions, the most common method involves the treatment of the parent halide MXn with the halide of a suitable univalent cation, e.g. an alkali-metal or substituted ammonium, phosphonium or arsonium species. Interaction may be conveniently brought about in the melt, in aqueous hydrohalic acid or in a non-aqueous solvent like acetonitrile, ethanol or chloroform327»343; in certain cases a non-aqueous solvent which is itself a halogenating agent, e.g. thionyl chloride348, has been favoured. The nature of the product is often strongly influenced by the relative proportions of the two halides, by the shape as well as the size of the cation, and by the reaction medium employed. In these procedures no deliberate attempt is made to vary the oxidation state of M. Complexes have also been made, however, via the reduction or oxidation, by chemical or electrolytic means, of solutions containing halide species, e.g. anodic dissolution in 6 M

► Ga 2 X6 2 ~

Ga

349

HX, Isolated as Me 4N+ salts Zn in HC1 solution

TcCle 2 " 2

ReOCls "

oxidation by air, HC1

^Tc2Cl83-

35

°

^ R e O C l 6 2 - 35i

Properties of Halides Table 21 provides a qualitative summary of some of the physical characteristics of halide systems MX n , emphasizing how these characteristics vary with the nature of the halide. In keeping with such variations, chemical behaviour also ranges from that attribut­ able to the halide ion X~, in various conditions of coordination, to the more specific properties diagnostic of the M-X bonds of molecular halides. Accordingly, at one extreme, the chemical properties of ionic halides are determined largely by changes of coulombic energy, through lattice or solvation energy terms, while, at the other, many of the reactions 346 K . Dehnicke, Angew. Chem., Internat. Edn. 4 (1965) 2 2 . F. Calderazzo, Halogen Chemistry (ed. V. Gutmann), Vol. 3, p. 383. Academic Press (1967). 348 D . M . A d a m s , J. Chatt, J. M . Davidson and J. Gerratt, / . Chem. Soc. (1963) 2189. 349 c . A . Evans and M . J. Taylor, Chem. Comm. (1969) 1201. 350 j . D . Eakins, D . G . Humphreys and C . E . Mellish, / . Chem. Soc. (1963) 6012. 351 R. Colton, Austral. J. Chem. 18 (1965) 435. 347

1258

CHLORINE. BROMINE, IODINE AND ASTATINE: A. J. DOWNS AND C. J. ADAMS

of molecular halides can be referred to the strength of the M-X bond, as expressed, for example, by its mean bond energy. To explore more closely the nature of the interaction between M and X and its dependence on the identities of M and X, reference is most commonly made to one or more of the following physical properties, which quantify various aspects of the M · · · X interaction329: the mean bond energy; the internuclear distance; the stretching force constant and related vibrational properties; esr, nmr, nqr and Mössbauer parameters; the ligand-field strength, interelectronic repulsion integrals and other terms derived from electronic spectra and magnetic properties; and dipole moments. Brief consideration is now given to these properties. 1. Mean bond energy, B(M-X). Defined by A// a t, the enthalpy of the process MXn(g) -> M(g)+«X(g), such that A// at = nB(M-X), the mean bond energy can be evalu­ ated more or less reliably for many halide molecules293»297. An alternative approach, appropriate to metal halides and to complex ions, involves the so-called "coordinate bond energy", which refers to the energy per bond for the heterolytic process MXrc(g) ->M^ + (g)+«X~(g); compilations of this energy term are to be found else­ where329. In either case, it should be noted, the energy required to break an individual bond may be quite different from the average value. Table 23 lists bond energy terms for most neutral halide molecules for which experi­ mental data are available. Spectroscopic methods have been used directly to determine the dissociation energy of some of the diatomic halide molecules352. More often, however, thermochemical methods furnish the heat of formation of gaseous MXn; with a knowledge of the heats of atomization of M and X, B(M-X) is then easily gained297»329»353. Despite the interpretative shortcomings of bond energy terms, some interesting trends may neverthe­ less be discerned in the results of Table 23. In the first place, for a given element M (M Φ F), B(M-X) diminishes in the order F > Cl > Br > I. Further, in relation to the typical non-metals M1? M 2 . . . of one of the vertical Groups IV-VII, the bond energy order is usually ^(M^X) < B(M2-X) > B(M3-X) > J?(M4-X) for a given halogen; by contrast, for the typical elements of Groups I—III, B(M-X) appears to diminish regularly as the atomic number of M increases. Moving from left to right across a given horizontal Period sees B(M-X) first rise, reach a maximum, and then fall to relatively low values for the interhalogens and noble-gas fluorides; this trend is modulated by a subsidiary minimum, which is reached with the completion of the d-shell at the end of each transition series. Unfortunately, the sparseness of accurate data precludes any meaningful generalization about the bond energies of analogous halides formed by metals of the different transition series, though there is some indication that, where the highest oxidation state is realized, the heaviest metals enjoy the largest bond energies. Where an element forms more than one halide MX n , it is usually the rule that the bond energy decreases as n increases. Compared with the halides formed by typical class a acceptors, e.g. Rb 1 or Sr n , halides of class b acceptors like Au1 or Hg n are notable, not only for their comparatively low bond energies, but also, significantly, for the comparatively small changes in bond energy that accom­ pany variations of X, e.g. 5(M-F) - 5(M-I) = 28, 29, 39 and 52 kcal for M = Au1, Hg n , Rb 1 and Sr n , respectively. 352 T. L. Cottrell, The Strengths of Chemical Bonds, 2nd edn., Butterworths, London (1958); A . G . Gaydon, Dissociation Energies, 3rd edn., Chapman and Hall, London (1968). 353 R . c . Feber, Los Alamos Report, U.S. At. Energy Comm. LA-3164 (1965).

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I N O ON t*·

vb pTt co

■«-» PH

2

1 >

Ö

Co *n NO m

HI

Tt

Ü 1 1

»n oo co i—i oo f - r - NO

HHP. r-
B

v o NO NO r o ON NO « n T t

NO 00 ON T t O 00 t ^ NO 1 *"^

1

OX)

B

1 HH

| PQ

ON CO t » NO

© oo

T-H

H

B

o

> > >>

O ON ON Poo vo m Tt

P-4

© τ-i B

00

T3 00 r-

ON

3 3

«e

00 T t «Λ N

HJ

33

Γ- OO O 00

1 PHOPPW

1

fcO

«HH

Superscript Roman numerals refer to the valence state of the element forming the halide. Data derived from: a F. D. Rossini, D. D. Wagman, W. H. Evans, S. Levine and I. Jaffe, Selected Values of Chemical Thermodynamic Properties, Circular 500, National Bureau of Standards, Washington (1952); National Bureau of Standards Technical Notes 270-1, 270-2, 270-3 and 270-4, U.S. Govt. Printing Office, Washington (1965-9). b T. L. Cottrell, The Strengths of Chemical Bonds, 2nd edn., pp. 270-289. Butterworths, London (1958). c R. C. Feber, Los Alamos Report LA-3164 (1965). d C. J. Cheetham and R. F. Barrow, Adv. High Temperature Chemistry, 1 (1967) 7; D. L. Hildenbrand, ibid. p. 193. e R. G. Pearson and R. J. Mawby, Halogen Chemistry (ed. V. Gutmann), Vol. 3, p. 55, Academic Press (1967). f D. A. Johnson, Some Thermodynamic Aspects of Inorganic Chemistry, Cambridge (1968). « Tables of Constants and Numerical Data, No. 17, Spectroscopic Data relative to Diatomic Molecules (general ed. B. Rosen), Pergamon (1970). h H. C. Ko, M. A. Greenbaum, M. Färber and C. C. Selph, / . Phys. Chem. 71 (1967) 254. 1 H. Schäfer, H. Bruderreck and B. Morcher, Z. anorg. Chem. 352 (1967) 122. J H. H. Rogers, M. T. Constantine, J. Quaglino, jun., H. E. Dubb and N. N. Ogimachi, / . Chem. and Eng. Data, 13 (1968) 307; W. R. Bisbee, J. V. Hamilton, J. M. Gerhauser and R. Rushworth, ibid. p. 382. k O. M. Uy, D. W. Muenow and J. L. Margrave, Trans. Faraday Soc. 65 (1969) 1296. 1 P. A. G. O'Hare and W. N. Hubbard, J. Phys. Chem. 69 (1965) 4358. m S. R. Gunn, / . Amer. Chem. Soc. 88 (1966) 5924. n B. Weinstock, E. E. Weaver and C. P. Knop, Inorg. Chem. 5 (1966) 2189. 0 D. Cubicciotti, Inorg. Chem. 7 (1968) 208, 211. P P. Gross and C. Hayman, Trans. Faraday Soc. 60 (1964) 45. * P. Gross, C. Hayman, D. L. Levi and G. L. Wilson, Trans. Faraday Soc. 58 (1962) 890.

1262

CHLORINE, BROMINE, IODINE AND ASTATINE: A. J. DOWNS AND C. J. ADAMS

Though it is tempting to relate such fluctuations of bond energy to changes either in the efficiency of orbital-overlap or in the availability of potential donor or acceptor orbitals, it must be appreciated that the intrinsic energies of the bonds are inevitably concealed by the general inaccessibility of information about electronic promotion energies separating the ground and valence states. In more practical terms, the bond energy sequences reflect the relative stabilities of halides with respect to thermal decomposition and to substitution reactions. Thus, whereas OF2 dissociates into its elements only at temperatures above 250°C, OCl2 is thermodynamically unstable at 25°C, exploding on heating or sparking, OBr2 begins to decompose above — 50°C, and OI2 is as yet unknown. Again, the dissocia­ tion of HI on mild heating contrasts with the explosive formation of HF and HCl from their elements at room temperature, and with the combination of H 2 and Br2 to give HBr in the presence of a catalyst at 200°C. The lower thermal stability of iodides has been exploited in the preparation of very pure elements, e.g. silicon, boron, titanium and thorium, by pyrolysis of the iodide on a hot wire293. 2. Bond length, r(M-X) 352 ' 354 » 355 . Studies of the microwave spectra, of rotational detail superimposed on electronic or vibrational spectra, or of the electron diffraction pat­ terns due to gaseous halide molecules MXn lead, in principle, to information about molecular dimensions. Determined for the most part by one or more of these methods, the M-X bond lengths of Table 24 refer exclusively to gaseous molecules; in the condensed phases intermolecular interactions may lead to significant changes in these lengths or even to an aggre­ gate in which the molecule ceases to be a recognizable unit. Where the M-X distance has been measured in a number of related molecules, e.g. MXW -mYm, a mean value is given, though, in common with bond energies, such distances are not strictly independent of the nature and number of the other substituents Y. For a given element M, r(M-X) increases in the order F < Cl < Br < I in a series of halides of the type MXW. Likewise, for a given halogen and a particular Group of typical elements, there is an attenuation of r(M-X) as the atomic number of M increases, though, in any horizontal Period, the normal trend with respect to atomic number is in the opposite sense. Evidently there exists no simple correlation between bond length and bond energy. Although short M-X bonds are usually characterized by high bond energies, with an inverse correlation between the length and energy, there are numerous series in which the trend appears to be reversed. Thus, of the diatomic halogen fluorides F2 exhibits at once the shortest internuclear distance and the lowest bond energy. Deviations of bond distance from the additivity principle (implicit in the designation of covalent radii for atoms) are presumed to reflect the influence of such variables as polarity, ττ-bonding or intermolecular interactions. Despite the appeals commonly made to bond-contraction as a sign of partial multiple bonding, as in S1CI4, it is difficult to unravel the separate contributions made by such bonding and by the polarity of the bond. The difficulty is exacerbated, moreover, by the fact, manifest in the data of Table 24, that distance is not a very sensitive function of bond character. 354

L. E. Sutton (ed.), Tables of Interatomic Distances and Configuration in Molecules and Ions, The Chemical Society, London (1958); Supplement (1965). 355 Tables of Constants and Numerical Data, No. 17, Spectroscopic Data relative to Diatomic Molecules (general ed. B. Rosen), Pergamon (1970).

1

2-34"

2-361

—

2-5021

2-7121

Br

I

2-270

2-7871 2-9451 3-1771

F

Cl Br I

1

— — —

2-076

1

2-4391 — —

2-6671 2-8211 3 048 1

Cl Br I

Rb

1-931

2-1721

F

K

2-18"

2-1981

2-3611

Cl

2-67" 2-82" 3 03"

2-20"

Sr

2-51" 2-67" 2-88"

210"

Ca

2-52"

1-77"

1-750

1-926

Mg

1-40" 1-75" 1-91" 2-10"

F

1-361 1-77" l-89 If 2-12If

1

Be

1

Na

1-564 2-0181 2-1701 2-3921

Li

0-9171 1-2751 1-4151 1-6091

1

F Cl Br I

F Cl Br I

H

Si

1

— — 2-44 m k

2-2021 2-3531 2-5751

2-43" 2-55" 2-78"

2-33^ 2-46™ 2-69^

— 2-325"1 2-51"1 2-67"1

Sb

Sn

2-4011 2-5431

2-7541

As

—

210v

1-712"1

2-47"1

2-20"1

l-58 2020VO

Vn

\2-124

2 .Oi0 3 " i / Z

l-563"

P Im

2-161"1 2-33"1 2-54"1

IV

1

1-36" 1-75"1

N

208IV 2-32IV 2-48IV

206"

1-985

Ge 1-68IV

1

In

l-88" Ik

1-7751

Ga

2-435IV

2-44" Ik

_

216 I V

—

2019 I V

l-591 l-561

ni

2-2961

l-63"

Ik

2-06" Ik

Al

IV

1-333 1-311™* 1-30" 1-767IV 1-741" 1 1-87" 1-938IV 2-10 IIIh 2-14IV

1

C

2-1301

1-654

1

1-262 1-7161 1-891 _

1

B

1.99vi

207 1 V

1-60 P 1-56 V I P

IV

Se

2-36" 2-33IV 2-51"

1-84

VI

Te

1 0 /

2-27IV 2190 V I

fl-638^ \ 1-771

2-24"

1-99"

1-635"P

S

1-42" l-693"i

O

F 1

1

1

2-3211 2-4851 2-6661

1-909

I

/1-7211" \ 1-810 2-1381 2-2811 2-4851

1756

le7W

Br

2-3211

2-1381

L ^ o ! fl-698 " \ 1-598 1-9881

Cl

1-418 1-6281 1-7561 1-9091

1628

TABLE 24. M-X BOND LENGTHS IN MOLECULAR HALIDES IN THE VAPOUR PHASE (A) a ~ e

1.97711s

J.94IVS l.890Vis

Xe

l-875" r

Kr

AT

Ne

He

4*

1

1·925 Ι 2·04™ — 2-47™ — 2-63™ — 2-80™

Y

1-7881 l - 9 1 m 2-231 —

Sc

2-32" 2-82" 2-99" 3-20"

Ba

2-22™ 2-60™ Rare 2-75™ Earths 2-98™

La

2-026 — — [—

Cl Br I

1F

F Cl Br I

3-315

2170

1

2.331VU 2.431VU

Hf

2.441VU

2-32IVu

Zr

2-185 Ivt 2-31 Ivt

Ti

v

l-86 2-27 v 2-44 v

Ta

l-88 v 2-28Vu 2-45Vu

Nb

1-71V 2-14 IV 2-12 vt 2-30 IV

V

1

2-084 2-4851 2-6181 2-8141

Tl

2-26 v l u

2-833VIV

W

l-826 v i u 2-23 IV 2-27 Vu 2-39IV

Mo

2-12 VIt

Cr

2-13" 2-46" 2-60" 2-81"

v

u

l-859 " 2-230 v " u

Re

Tc

1-70 I 1-724 V "

Mn

— 2-43 IV — —

Pb

l-831

Os

Ru

Fe

2-48™ 2-63™

Bi

viv

l-830

Ir viv

l-83 v l v

Rh

Co

Po

l-83

Pt

Pd

Ni

viv

At

Au

I.991 2-2811 2-3921 2-5441

Ag

1-7491 2050 1 2-1731 2-3351

Cu

2-05" 2-29" 2-41" 2-59"

Hg

1-97" 2-235" 2-37" 2-55"

Cd

1-81" 2-05" 2-21" 2-38"

Zn

Rn

Superscript Roman numerals refer to valence state of the element forming the halide. Equilibrium bond distances re are given wherever the necessary information is available. Distances are derived from the following sources: a T. L. Cottrell, The Strengths of Chemical Bonds, 2nd edn., pp. 270-289. Butterworths, London (1958). b Tables of Interatomic Distances and Configuration in Molecules and Ions (ed. L. E. Sutton), The Chemical Society, London (1958); Supplement (1965). c C. J. Cheetham and R. F. Barrow, Adv. High Temperature Chemistry, 1 (1967) 7. d L. V. Vilkov, N. G. Rambidi and V. P. Spiridonov, /. Struct. Chem. 8 (1967) 715. e Tables of Constants and Numerical Data, No. 17, Spectroscopic Data relative to Diatomic Molecules (general ed. B. Rosen), Pergamon (1970). f A. Snelson, /. Phys. Chem. 11 (1968) 250. β K. Kuchitsu and S. Konaka, /. Chem. Phys. 45 (1966) 4342.

I

F Cl Br

F Cl Br I

2-3451 2-9061 3-0721 1

Cs

TABLE 24 (cont.)

1

A. G. Massey, Adv. Inorg. Chem. Radiochem. 10 (1967) 1. F. X. Powell and D. R. Lide, jun., / . Chem. Phys. 45 (1966) 1067. i B. Beagley, A. H. Clark and T. G. Hewitt, / . Chem. Soc. (A) (1968) 658. k Refers to monomer unit, see ref. d. 1 V. M. Rao, R. F. Curl, jun., P. L. Timms and J. L. Margrave, / . Chem. Phys. 43 (1965) 2557. m E. Hirota and Y. Morino, / . Mol. Spectroscopy, 33 (1970) 460. n K. W. Hansen and L. S. Bartell, Inorg. Chem. 4 (1965) 1775, 1777; S. B. Pierce and C. D. Cornwell, / . Chem. Phys. 48 (1968) 2118. 0 W. J. Adams and L. S. Bartell, / . Mol. Structure, 8 (1971) 23. p H. L. Roberts, Inorganic Sulphur Chemistry (ed. G. Nickless), p. 419, Elsevier (1968); Essays in Structural Chemistry (ed. A. J. Downs, D. A. Long and L. A. K. Staveley), p. 457, Macmillan (1971). * I. C. Bowater, R. D. Brown and F. R. Burden, / . Mol. Spectroscopy, 23 (1967) 272; ibid. 28 (1968) 454, 461. r C. Murchison, S. Reichman, D . Anderson, J. Overend and F. Schreiner, / . Amer. Chem. Soc. 90 (1968) 5690. 8 J. H. Holloway, Noble Gas Chemistry, Methuen, London (1968); R. M. Gavin, jun., and L. S. Bartell,/. Chem.Phys. 48 (1968) 2460, 2466; S. Reichman and F. Schreiner, / . Chem. Phys. 51 (1969) 2355. % R. Colton and J. H. Canterford, Halides of the First Row Transition Metals, Wiley-Interscience (1969). u J. H. Canterford and R. Colton, Halides of the Second and Third Row Transition Metals, Wiley-Interscience (1968). v M. Kimura, V. Schomaker, D. W. Smith and B. Weinstock, / . Chem. Phys. 48 (1968) 4001; H. Kim, P. A. Souder and H. H. Claassen, / . Mol. Spectroscopy, 26 (1968) 46.

h

1266

CHLORINE, BROMINE, IODINE AND ASTATINE: A. J. DOWNS AND C. J. ADAMS

3. Vibrational properties: force constants. Various compilations 345 ' 352 » 355-359 bear witness to the research effort expended on investigations of the vibrational properties of simple and complex halides. Vibrational frequencies have been derived from the vibrational progressions observed in the electronic spectra of diatomic halides352»355 or from the infrared and Raman spectra of more complicated species345»356 ~359. Such studies have communicated much valuable information concerning the structure and bonding properties of halide systems. In a diatomic molecule MX, the frequency of the M-X stretching vibration is determined by two factors, the bond-stretching force constant and the masses of the two atoms. Hence, from the observed frequencies one may readily determine bond-stretching force constants, which are measures of the resistance of the M-X bond to deformation. In contrast with the bond energy, which refers to the process of separation of the M and X atoms to large dis­ tances, the force constant gives a measure of bond strength at internuclear distances close to the equilibrium value in terms of the curvature of the potential energy surface. Repre­ sentative force constants for some diatomic halides are contained in Table 25. In polyatomic systems, the situation is complicated by interactions between formally non-bonded atoms and by the fact that the normal modes involve, not only bond-stretching, but also angular and torsional deformations. Detailed normal coordinate analyses leading to realistic force constants have generally been feasible only for halide molecules or complex ions of relatively high symmetry; even then, there is a need for additional information concerning, typically, isotopically substituted species or Coriolis coupling constants, if the secular equation is to be solved without undue simplification of the molecular force field. If most of the energy of a normal mode is localized in the deformation of a given bond or molecular unit, the frequency of the mode is then characteristic of that bond or unit. The condition whereby a normal mode may be considered as a "group vibration" implies that the frequency of the motion is well separated from others of the same symmetry class. Essentially localized vibra­ tions are assumed for the identification of M-X "stretching frequencies". A popular basis for qualitative analysis and, less reliably, for inferences about chemical bonding, the range and magnitude of such frequencies have been discussed in some detail345»358»359; the results have been tabulated or used to construct correlation charts 345 » 356-359 . However, "group frequency" arguments affecting bonds in which the heavier halogen atoms are en­ gaged are at best approximate. For species of high symmetry, e.g. tetrahedral MX4 or octahedral MX6, there exist unique totally symmetric modes localized in the M-X bond coordinates, the frequencies of which can be meaningfully correlated with the M-X stretch­ ing force constant360. In more complex or less symmetrical units, e.g. [MoöClgCle]2 ~ 361, it is more often the rule that the normal modes involve extensive mixing of the various bond vibrations. Consideration of the mass of data available discloses that the stretching frequencies 356 K. Nakamoto, InfraredSpectra of Inorganic and Coordination Compounds, 2nd edn., Wiley-Interscience, New York (1970). 357 H. Siebert, Anwendungen der Schwingungsspektroskopie in der Anorganischen Chemie, SpringerVerlag (1966). 358 D. M. Adams, Metal-Ligand and Related Vibrations, Edward Arnold, London (1967). 3 59 R. J. H. Clark, Halogen Chemistry (ed. V. Gutmann), Vol. 3, p. 85. Academic Press (1967). 36 0 L. A. Woodward, Trans. Faraday Soc. 54 (1958) 1271. 361 M. J. Ware, Physical Methods in Advanced Inorganic Chemistry (ed. H. A. O. Hill and P. Day), p. 241. Interscience (1968).

1-37 0-86 0-70 0-53

KF KC1 KBr KI

b

CaF CaCl CaBr Cal

2-62 1-51 1-27 102

MgF 3-25 MgCl 1-82 MgBr 1-51 Mgl - 1 1 6

MnF 3-18 MnCl 1-82 MnBr 1-57 Mnl - 1 - 3 0 CuF CuCl CuBr Cul

3-34 2-31 2 05 1-74

GaF GaCl GaBr Gal

A1F A1C1 AlBr All 3-40 1-83 1-52 1-24

4-22 209 1-70 1-31 GeF GeCl GeBr Gel

SiF SiCl SiBr Sil 3-92 2-32 1-97 1*65

4-91 2-64 2-22 1-76

— —

— —

4-98 3-24

AsCl 2-78

PF PCI

FOR SOME DIATOMIC HALIDE MOLECULES*1*

— — —

— — —

4-48 3-23 2-80 2-39

BrF 4 0 9 BrCl 2-80 Br2 2-48 BrI 2 07

C1F Cl 2 ClBr C1I

T. L. Cottrell, The Strengths of Chemical Bonds, 2nd edn., pp. 270-289. Butterworths, London (1958). Tables of Constants and Numerical Data, No. 17, Spectroscopic Data relative to Diatomic Molecules (general ed. B. Rosen), Pergamon (1970).

1-76 110 0-96 0-76

NaF NaCl NaBr Nal

a

9-66 516 4-11 3-14

HF HC1 HBr HI

TABLE 25. STRETCHING FORCE CONSTANTS, ke (mdyne A - 1 )»

1268

CHLORINE, BROMINE, IODINE AND ASTATINE! A. J. DOWNS AND C. J. ADAMS

v(MX) or, more precisely, the force constants k(MX) of bonds to halogen atoms are func­ tions of the following variables358»359: The ionic character of the bond The general pattern (see Table 25) is that v(MX) and fc(MX) diminish as the polarity of the M-X bond increases, reflecting the variation in form of the potential energy curve as the nature of the bond changes. The nature of M and X It appears to be a general rule that, for a given element M, v(MX) and k(MX) decrease in the order F > Cl > Br > I; in general, too, an inverse correlation with the atomic number of M is found in vertical Groups of the typical elements, but for halides formed by analogous elements of the Ad and 5d transition series, e.g. MoF 6 and WF 6 , PdCl62~ and PtCl62 ~, CdCl2 and HgCl2, there is a significant reversal of this trend. The vibrational properties of transition-metal-halogen bonds also reflect variations of ligand-field stabiliza­ tion energy and of the bonding or anti-bonding character of the ^/-electron shell358»359. Stereochemistry and coordination number The geometry of a halide aggregate has a profound influence on its vibrational properties in terms, not only of selection rules, but also of the frequencies of the vibrational modes. The general rule has been propounded358»359 that v(MX) decreases with increasing co­ ordination number of M or X: hence, bridging halogen units, as in Al2Cl6, are normally characterized by lower v(MX) and A:(MX) than are terminal M-X bonds. Oxidation state and charge An increase in the oxidation state of M is normally accompanied by an increase in v(MX) and k(MX). If M is varied in an isoelectronic series, e.g. AsCl 4 + , GeCl4, GaCl 4 -, ZnCl42 ~, the same trend is observed, namely, diminution in v(MX) and k(MX) as the nega­ tive charge borne by the aggregate increases, an effect which has been related to the radial electronic distribution of M 360 . The presence of other ligands There is ample evidence that the vibrational properties of M-X bonds are sensitive to the nature and, in some cases, the location of other ligands L in a mixed aggregate such as SnX w L 4 -« or PtX 2 L 2 . Influences such as the electronegativity of L or potential ττ-bonding between M and L have been weighed in these circumstances358»359. Environmental effects The aggregation of molecules or ions in a crystal is always liable to alter the vibrational properties of these units. Contributory to such changes are the effective symmetry of the molecule or ion, the site of which usually belongs to a symmetry group different from that of the isolated unit, and, in the case of complex ions, the nature of the counter-ion, which may have a marked frequency- or even structure-determining influence. The behaviour of halide species in solution is inevitably dependent on the nature of the solvent, though there have been few systematic studies of the effects of solvent on the vibrations of simple and complex halides358»359.

GENERAL PROPERTIES OF HALIDES

1269

4. Esr studies 329 ' 362-368 . Measurements of the electron spin resonance (esr) spectra of paramagnetic transition-metal halides afford information about three parameters, the g-factor, spin-orbit coupling constant λ, and nuclear hyperfine splitting constant A, all of which can be related to the degree of electron-delocalization. The principles are well exemplified by the classical studies of the IrCl62 - and IrBr62 - ions in magnetically dilute mixed crystals362. The esr spectra show hyperfine structure lines arising from interaction of the magnetic electrons both with the Ir nucleus (/ = 3/2) and with the Cl or Br nucleus (/ = 3/2). From the A- and ^-values it has been deduced that the "hole" in the formally non-bonding t2g metal orbitals spends 3-5% of its time on each of the six halogen atoms, and that the hole and five t2g electrons actually occupy molecular orbitals derived from combinations of the metal 5dxy-, dX2- and dy2-orbitals and suitable /Vtype orbitals of the halogen ligands. Other examples of paramagnetic halides studied in this way include the complex ions CoX64 -, for which the extent of delocalization increases in the sequence F < Cl < Br < I, and fluorides and chlorides of the lanthanides, in which the delocalization of the 4/-electrons appears to be strictly limited364'366»367. Tabulations of esr data relating to bonding in halides are to be found elsewhere329'363'364»367. 5. Nmr studies. The interactions between ligand nuclei and the magnetic electrons of transition-metal ions, which lead to the hyperfine structure observed in the esr spectra, also result in large shifts in the magnetic resonance of the ligand nucleus. Studied in a number of transition-metal fluorides329»364'369, these shifts have led to estimates of the fraction of unpaired electrons in fluorine 2s-, 2ρσ- and 2/7^-orbitals, the results being in good agreement with those deduced from esr measurements. Apart from studies of crystalline FeCl2 369 and of some solutions containing chloride in the presence of a transition-metal ion, however, the influence of electron paramagnetism on the resonances due to the heavier halogen nuclei has so far received little attention. Measurements of the magnetic resonance of the heavier halogen nuclei in molecular halides have been virtually restricted to 35C1, though a few relaxation times for other nuclei have been determined indirectly by chemical exchange experiments307. Table 26 lists the 35 C1 chemical shifts and relaxation times (T2) of some representative chloride species. Interpretation of the chemical shifts is hampered by the difficulty of rigorous calculation, which requires knowledge of the ground and excited-state wavefunctions. A relatively simple practical equation given by Saika and Slichter for the shift of the 19 F resonance in the F 2 molecule relative to the free ion370 has provided a basis for the discussion of the 35 C1 shifts in molecular chlorides, as for example in the case of the CI2 molecule307»371. 362 j . O w e n a n d K . W . H . Stevens, Nature, 171 (1953) 8 3 6 ; J . H . E . Griffiths, J . O w e n a n d I . M . W a r d , Proc. Roy. Soc. A219 (1953) 5 2 6 ; J . H . E . Griffiths a n d J . O w e n , Proc. Roy. Soc. A226 (1954) 96. 36 3 B . R . M c G a r v e y , Transition Metal Chemistry ( e d . R . L . C a r l i n ) , V o l . 3 , p . 89. E d w a r d A r n o l d a n d Marcel D e k k e r (1966). 364 j . O w e n a n d J . H . M . T h o r n l e y , Rep. Progr. Phys., London, 2 9 (1966) 675. 365 A . C a r r i n g t o n a n d A . D . M c L a c h l a n , Introduction to Magnetic Resonance, H a r p e r a n d R o w (1967); A . C a r r i n g t o n a n d H . C . Longuet-Higgins, Quart. Rev. Chem. Soc. 14 (1960) 4 2 7 . 366 A . Abragam a n d B . Bleaney, Electron Paramagnetic Resonance of Transition Ions, Clarendon Press, Oxford (1970). 367 B . A . G o o d m a n a n d J . B . R a y n o r , Adv. Inorg. Chem. Radiochem. 1 3 (1970) 135. 368 L . Shields, / . Chem. Soc. (A) (1971) 1048 a n d references cited therein. 369 D . R . E a t o n a n d W . D . Phillips, Adv. Magnetic Resonance, 1 (1965) 103. 370 A . Saika a n d C . P . Slichter, J. Chem. Phys. 2 2 (1954) 2 6 . 371 C. Hall, D . W. Kydon, R. E. Richards and R. R. Sharp, Mol. Phys. 18 (1970) 711.

1270

CHLORINE, BROMINE, IODINE AND ASTATINE: A. J. DOWNS AND C. J. ADAMS

TABLE 26.

35

C1 CHEMICAL SHIFTS, RELAXATION TIMES AND QUADRUPOLE COUPLING CONSTANTS FOR SOME CHLORINE-CONTAINING SPECIES*

Species

Temperature (°K)

Chemical shift* (ppm)

C10 3 -(aq) C10 4 _ (aq) TiCl4 VOCI3 SO2CI2 Cr0 2 Cl 2 CCI4 S2CI2 POCI3 CHCI3 Cl2 PCb CH2C12 SiCl4 GeCl4 Cl 3 -(aq)

Room 299 299 299 299 299 299 299 299 Room 298 299 Room 299 299 298

-1050 -946 -840 -791 -760 -603 -500 -480 -430 -410 -370 -370 -220 -174 -170 150

AsCl3 SnCl4 CH3CI Cl-(aq), dil soln*

299 299 Room Room

-150 -120 -40 0

Relaxation time, r2(msec) 0024 51 0-40 0161 0030 0107 0025 0 022 0 027 0025 0031 0060 0042 0079 0041 0003t 0015J 0023 0022 0099

—

Quadrupole coupling constant, e2Qqlh(MHz)b

—

12-2 23-1 75-4 31-4 81-2 71-6 57-9 76-7 109 52-2 720 40-8 51-4 116f 51ί 50-3 48-2 681

—

* External reference, t Central atom, t Terminal atom. * C. Hall, Quart. Rev. Chem. Soc. 25 (1971) 87. to E. A. C. Lücken, Nuclear Quadrupole Coupling Constants, Academic Press, London and New York (1969).

With 35C1, as with 19 F, the nuclear magnetic shielding appears to be dominated by the para­ magnetic contribution370 ~372, which has been related by the expression

— KSKKiX, to the mean excitation energy ΔΕ, effectively expressing the accessibility of electronic states above the ground state, and to the expectation value of r - 3 for the valence /^-electron. The periodic dependence of the range of chemical shifts on the atomic number of the nucleus has been shown372 to follow the pattern dictated by the term wp ; for the halogens, these ranges are (in ppm): F, 625; Cl, 1000; Br, ~ 1650. More recently, Johnson, Hunt and Dodgen373 have sought to express the 35C1 chemical shifts in terms of the quadrupole coupling constant e2Qq, an accurately known, independent experimental parameter, obtaining "par

*(e2Qqlh)AE-Hllr*>nP

Although the mean excitation energy is not clearly defined, transition-metal halides which 372 c . J. Jameson and H . S. Gutowsky, / . Chem. Phys. 4 0 (1964) 1714. 373 K. J. Johnson, J. P. Hunt and H . W. D o d g e n , / . Chem. Phys. 51 (1969) 4 4 9 3 .

1271

GENERAL PROPERTIES OF HALIDES

have relatively low-lying excited states, but rather small quadrupole coupling constants, are mostly characterized by large chemical shifts. Furthermore, there is a general trend in the direction of larger shifts in compounds with large quadrupole coupling constants. However, there are also distinct anomalies, e.g. the wide variation of chemical shift for the series CH3C1, CH2C12, CHC13, CC14. There exist many data relating to the chemical shifts 374-379 and coupling conStantS374>379>380

of

nuclei

SUCh as

Ή374,

ΠΒ 374 > 37 8,

13C374.377,

14^74,376,

19F3749

3ip374,3755 o r I95pt379 i n halogen-containing molecules. In the interpretation of chemical shifts, the shielding of a nucleus is conveniently expressed as the algebraic sum of the local diamagnetic (ad) and paramagnetic (σ^) contributions. For nuclei other than hydrogen, variations in shielding are largely dictated by variations in σρ, which, being sensitive to changes in the wavefunction of the valence electrons, reflects the bonding properties of the resonant atom. Hence the 195Pt chemical shifts of the platinum hydrides trans[Pt(PEt3)2HX] have been correlated with changes in the σ- and π-character of the Pt-X bond379. However, recent analyses376»377 of carbon and nitrogen compounds have revealed that changes in σα can be quite large and that errors of interpretation can follow the neglect of these terms. In general, the complexity of factors which influence the magnetic shielding and interaction of nuclei tends to frustrate attempts to reach unambiguous conclusions about chemical bonding, though there are to be found numerous correlations, commonly of an empirical nature, with properties such as bond length, bond angle, hybridization and electronegativity374.

6. Nqr studies329»365»381 ~386. The energy of a non-spherical atomic nucleus in an inhomogeneous electric field varies according to the orientation of the nucleus about some fixed axis. The interaction between the nucleus and field gives rise to so-called nuclear quadrupole coupling; an inhomogeneous electricfieldis provided, for example, by the incompletely filled /?-shell found in halogen atoms. In a nuclear quadrupole resonance (nqr) experiment, radiation in the radiofrequency region is employed to effect transitions among the various orientations of a quadrupolar nucleus in an asymmetric intramolecular field; the frequencies of the transitions depend upon both thefieldgradient q produced by the valence electrons and 374 J. A. Pople, W. G. Schneider and H. J. Bernstein, High-resolution Nuclear Magnetic Resonance, McGraw-Hill (1959); J. W. Emsley, J. Feeney and L. H. SutclifFe, High Resolution Nuclear Magnetic Resonance Spectroscopy, Pergamon (1965-6). 375 M . M. Crutchfield, C. H. Dungan, J. H. Letcher, V. Mark and J. R. Van Wazer, P3i Nuclear Magnetic Resonance, Interscience (1967). 376 R . Grinter and J. Mason, / . Chem. Soc. (A) (1970) 2196. 377 j . Mason, / . Chem. Soc. (A) (1971) 1038. 378 See for example H. Nöth and H. Vahrenkamp, Ber. 99 (1966) 1049; M. F. Lappert, M. R. Litzow, J. B. Pedley and A. Tweedale, / . Chem. Soc. (A) (1971) 2426. 379 R . R. Dean and J. C. Green, / . Chem. Soc. (A) (1968) 3047. 380 w . McFarlane, Quart. Rev. Chem. Soc. 23 (1969) 187. 381 W. J. Orville-Thomas, Quart. Rev. Chem. Soc. 11 (1957) 162. 382 x. p. Das and E. L. Hahn, Nuclear Quadrupole Resonance Spectroscopy, Academic Press, New York (1958). 383 M. Kubo and D. Nakamura, Adv. Inorg. Chem. Radiochem. 8 (1966) 257. 384 H . Sillescu, Physical Methods in Advanced Inorganic Chemistry (ed. H. A. O. Hill and P. Day), p. 434. Interscience (1968). 385 E . A. C. Lücken, Nuclear Quadrupole Coupling Constants, Academic Press (1969). 386 E. A. C. Lücken, Structure and Bonding, 6 (1969) 1.

1272

CHLORINE, BROMINE, IODINE AND ASTATINE: A. J. DOWNS AND C. J. ADAMS

the quadrupole moment eQ of the nucleus. Structural information about a compound can be obtained by considering how different structural and electronic effects influence q, which is an index to the asymmetry of the electron environment. The direct measurement of quadrupole coupling by absorption of radiofrequency radiation is restricted to the crystalline phase, where the axes of q are fixed in space; quadrupole coupling constants for species in the gas phase are most easily derived from the fine structure in the pure-rotational (micro­ wave) spectrum. Nuclear quadrupole coupling in solids has also been extensively investi­ gated by the following techniques: nuclear magnetic resonance, electron spin resonance, electron-nuclear double resonance (ENDOR) and Mössbauer spectroscopy. The naturally occurring quadrupolar nuclei 35C1, 37C1, 79Br, 81Br and 127I have been a mainstay of nqr experiments, many of which have been directed to exploring the properties of halogen-bearing bonds. An analysis of the factors affecting the field gradient experienced by a halogen atom in a molecule indicates that the major contribution arises from the way in which the p-orbitals are occupied by the valence electrons; by contrast, electron density in the spherically symmetrical ^-orbital does not give rise to a field gradient. At one extreme, represented by the free halogen atom, the /^-electron "hole" produces a large field gradient, e.g. e2Qq = —109-746 MHz for 35C1: at the other, represented by the free halide ion, the spherical symmetry of the closed electron shell produces a vanishing field gradient at the nucleus and e2Qq = 0. The measured nqr frequencies therefore give a relatively direct indication of the ionic character of a halide species (see Table 27), though considerable difficulties arise when a more quantitative treatment is attempted. Most interpretations have been based on the approximate valence-bond approach pioneered by Townes and Dailey387, which leads to the following general equation relating e2Qqexp, the measured quadrupole coupling constant for the halogen atoms in a halide species MXn, to e2QqQit, the corresponding constant for the isolated atom, by388 e2Ö<7exP = [l-s-d-i(l-s-d)]e2Qq&t

(13)

Here s and d represent the fractions of s- and ^-character in the σ-bonding orbitals of the halogen atom, and / denotes the degree of ionic character in the M-X bond. A modified form of the equation, namely, e2Qqexp = [l-s+d-i-7r]e2Qq&t

(14)

has also been employed389 to give an account of the extent of ττ-interaction in M-X bonds. In view of the discrepancy between the number of measurable and unknown parameters, a rigorous solution is not possible without additional information. Whereas a σ-bond to a halogen should produce an axially symmetric field gradient, ττ-bonding leads to asymmetry, the extent of such bonding being related to the asymmetry parameter η. For example, this has formed the basis of a detailed study of^bonding in halides of the Group IV elements390, and also affords cogent evidence of extensive ττ-bonding in the boron trihalide mole­ cules385. 387 c . H . Townes and B. P. D a i l e y , / . Chem. Phys. 17 (1949) 782; ibid. 2 3 (1955) 118;

B.

P.

Dailey,

Discuss. Faraday Soc. 19 (1955) 255. 388 R. S. Drago, Physical Methods in Inorganic Chemistry, Reinhold, New York (1965). 339 B. P. Dailey, / . Chem. Phys. 33 (1960) 1641. 390 R. Bersohn, / . Chem. Phys. 22 (1954) 2078; J. D. Graybeal and P. J. Green, / . Phys. Chem. 73 (1969) 2948.

F35C1 C1127I D79ßr Cl?9Br D127I 35d 2

1000 0-968 0-911 0-867 0-831 0-329

004 3 58 259-87 15-8 10890

K35C1 Cs35Cl Na?9Br Nai27l T135C1 F79Br

Ionic character* 0-78 (0-43) 0-72 (0-45) 0-62 (0-47) 0-55 (0-47) 0-44 0-43 0-66 0-63 0-56 0-68

Quadrupole coupling constant, e*Qqlh (MHz)

20-4 27-9 33-8 41-7 520 53-5 31-5 34-5 41-2 30-3 346 271

—

406 411 261

— —

231

Quadrupole coupling constant, e2Qqjh (MHz)

X = 79ßr

(0-39)

— —

0-47 0-58

—

0-38 0-37 0-60

0-66

Ionic character*

-2292-712.

X = 1271

1020

— —

0-48

— —

—

0-55

—

884

(0-32)

— —

0-30

0-57 1360

822

— —

Ionic character*

0-259 0-229 0186 0110 0065 0

Ionic character

Quadrupole coupling constant, e*Qq\h (MHz)

1460 2944 533 876-8 1823-38 108-95

Quadrupole coupling constant, e*Qq\h (MHz)

* The numbers in brackets stand for the total ionic character where ττ-bonding has been considered. a C. H. Townes and B. P. Dailey, / . Chem. Phys. 17 (1949) 782; ibid. 23 (1955) 118. b H. Sillescu, Physical Methods in Advanced Inorganic Chemistry (ed. H. A. O. Hill and P. Day), p. 456. Interscience (1968).

K2IrX6 K2PtXö K2PdXö K2SnXö (NH4)2PbX6 K2SeXö K2TeXö

K 2 OSXÖ

K2WX6 K2ReXö

Hexahalide

X = 35C1

(b) Complex halides*

eZQqlh atom (MHz): 35Q, -109-746; 37Q, -86-510; 79ßr, 769-756; sißr, 643032; mi,

Halide

Ionic character

Quadrupole coupling constant, e^Qq/h (MHz)

Halide

(a) Binary halides*

TABLE 27. NUCLEAR QUADRUPOLE COUPLING CONSTANTS AND IONIC CHARACTER IN SOME HALIDE SPECIES

1274

CHLORINE, BROMINE, IODINE AND ASTATINE: A. J. DOWNS AND C. J. ADAMS

By assuming the d- and ^-contributions to be unimportant and the extent of ^-hybridi­ zation to be 15%, Townes and Dailey387 have calculated the ionic character i for a large number of diatomic halides (see, for example, Table 27). Alternatively, the experimental quadrupole coupling constants have been used to calculate the ^-character of halogen bonds, the ionic character being defined through its relation to valence-state energies391. That the interpretation of the results for polyatomic halides385 has been less successful is indicated by the divergence of opinion expressed by various authors regarding the estimates of the quantities /, π, d and s in equation (14). A recent analysis392 suggests that there is a distinct change in the ^-character of the chlorine σ-orbital when chlorides pass from the solid to the gaseous state, and that the ^-character depends upon the atom to which the chlorine is bonded. However, the general conclusion, here and elsewhere386, that in very few cases do the results provide clear evidence of outer J-orbital participation, e.g. in Si-X, P-X or S-X bonds, must be tempered by the recognition that approximate methods are necessarily involved in the analysis of the experimental data. The problem of ionic character in transition-metal halides and hexahalo-complexes of heavier atoms has also been investigated by direct measurements of the nqr spectra383 ~385; some of the results are reproduced in Table 27. In some of these systems the relative contributions of σ- and ττbonding have been evaluated, but, in view of the burden of assumptions made, the conclu­ sions must be regarded as open to doubt327»384'388. A more general treatment of quadru­ pole coupling constants in terms of a molecular-orbital approach393 may serve to remove some of the uncertainties, but radical progress must be inhibited by the paucity of observ­ able parameters. 7. Mössbauer spectra 394-397 . An alternative source of information about nuclear quadrupole coupling, Mössbauer spectroscopy is based on the phenomenon of recoilless emission and resonance absorption of low energy y-radiation394. Chemical applications of this technique depend on hyperfine interactions between the nuclear energy levels and the extranuclear electrons, which give rise (i) to nuclear isomer shifts δ, (ii) to quadrupole splittings, and (iii) to magnetic hyperfine Zeeman splittings. To achieve recoilless condi­ tions for the nuclei under examination, studies of Mössbauer, as of nqr, spectra are essentially confined to solids. The quadrupole splitting pattern of a Mössbauer spectrum gives direct access to the quadrupole coupling constant e2Qq, and if / > 3/2 for one of the nuclear states implicated in the y-transition, the pattern may also disclose the asymmetry parameter η. The interpretation of such results in terms of site symmetries and field gra­ dients within a crystal and of the imbalance of p- and d-electrons then follows the pattern outlined in the previous section; hence, features such as ionic character, ^-hybridization and 7r-bonding have been analysed. 391 M . A . W h i t e h e a d a n d H . H . Jaffe, Theoret, Chim. Acta, 1 (1963) 209. 392 M . K a p l a n s k y a n d M . A . W h i t e h e a d , Mol. Phys. 16 (1969) 4 8 1 . 393 F . A . C o t t o n a n d C . B . H a r r i s , Proc. Nat. Acad. Sei. U.S.A. 56 (1966) 12. 394 E . Fluck, Adv. Inorg. Chem. Radiochem. 6 (1964) 4 3 3 ; N . N . G r e e n w o o d , Chem. in Britain, 3 (1967) 56; V. I. Goldanskii, Angew. Chem., Internat. Edn. 6 (1967) 830; R. H . Herber, Progress in Inorganic Chem­ istry, 8 (1967) 1; J. D a n o n , Physical Methods in Advanced Inorganic Chemistry (ed. H . A . O. Hill and P. D a y ) , p . 380, Interscience (1968); V. I. Goldanskii a n d R . H . H e r b e r (eds), Chemical Applications of Mössbauer Spectroscopy, A c a d e m i c Press (1968). 395 M . P a s t e r n a k , Symposia of the Faraday Soc. 1 (1967) 119. 396 D . W . Hafemeister, Advances in Chemistry Series, 68 (1967) 126. 397 G. J. Perlow, Chemical Applications of Mössbauer Spectroscopy (ed. V. I. Goldanskii and R. H . H e r b e r ) , p . 377. A c a d e m i c Press (1968).

1275

GENERAL PROPERTIES OF HALIDES

Of particular interest is the isomer shift parameter δ, which is unique to the Mössbauer experiment. This derives from the fact that an ^-electron has a finite probability of being found within the nuclear volume. Since the nucleus alters its radius r by a small amount 8r during a y-decay, there is a change in the fraction of the electron within the nucleus, with a concomitant change in electrostatic energy. If the source and absorber in a Mössbauer experiment are chemically distinct, first-order perturbation theory gives the shift in the resonance line as δ =4Lze2r2~ [W0)| 2 absorber - W0)|2source]

(15)

where Z is the nuclear charge, e the electronic charge and \Φ(0)\2 the total 5-electron density at the nucleus. The isomer shift is therefore the product of a nuclear term, which is a physi­ cal constant for a given isotope and can be determined, and a chemical term relating to the change in ^-electron density. By setting up suitable calibration scales, it is possible to moni­ tor the ^-electron density at a particular nucleus in a series of compounds, e.g. halide derivatives of 57 Fe, 119Sn and 197Au. As a measure of ^-orbital occupancy, the isomer shift depends upon the valence state and bonding orbitals of the atom, upon the shielding effects of /?-, d- or/-electrons, upon the ionic-covalent balance of interactions involving the atom, and upon the donor or acceptor capacity of neighbouring atoms through σ- or ττ-bonding. Of the halogen nuclei, only the two isotopes of iodine 127I and 129I are suitable for Mössbauer experiments (see p. 1158) 395-397 . Experimental measurements affecting these

_jlf a ii i I U -5

i -4

I L -3

I

Ιϊι i -2

i -1

ULLI

' n Hi' l' ill 0 +1

1_|

ι +2

I

I +3

I

d +4

I +5

FIG. 22. *29i isomer shifts (measured with respect to a ZnTe source) of iodine compounds.

isotopes have been focused, inter alia, on the following compounds: alkali-metal iodides396, CHnU-n (n = 0-3)398, MI 4 (M = C, Si, Ge or Sn)39*, IX (X = Cl, Br, I or C N ) 3 ^ * ^ IF 5 , IF 7 and related ions400, I 2 C1 6 399 , and derivatives of polyhalide, iodate and periodate ions 395-397 . Figure 22 illustrates the range of isomer shifts for 129 I; the spectra of many compounds also exhibit quadrupole splitting, whence the parameters e2Qq and η have been derived. Through their dependence on the 5.y-electron density at the nucleus, the isomer 398 B . S. Ehrlich and M . Kaplan, / . Chem. Phys. 5 0 (1969) 2 0 4 1 . 399 M . Pasternak and T. Sonnino, / . Chem. Phys. 4 8 (1968) 1997, 2009. 400 S. Bukshpan, C . Goldstein, J. Soriano and J. Shamir, / . Chem. Phys. 51 (1969) 3976; S. Bukshpan, J. Soriano and J. Shamir, Chem. Phys. Letters, 4 (1969) 241.

1276

CHLORINE, BROMINE, IODINE AND ASTATINE: A. J. DOWNS AND C. J. ADAMS

shifts in different iodine compounds provide an index to the relative participation of the 5s- and 5/?-electrons in the bonds formed by the iodine atom. In alkali-metal iodides, the isomer shifts are linearly dependent on the number ofp "holes" in the closed-shell con­ figuration 5s25p6, calculated independently from nmr and nqr data 396 ; this result has been used to calibrate the isomer shift in terms of the relative change in 5s density. Of the factors influencing the 5s density—electrostatic interactions, covalency, electronic shielding and overlap deformation of the free-ion orbitals—it appears that shielding and overlap make the major contributions. Where only /^-electrons are involved in the bonding, the removal of ^-electron density for bonding decreases the shielding of the ^-electrons and therefore increases the ^-electron density at the nucleus, with the result that the isomer shift relative to I " is positive (see Fig. 22). To a first approximation, the isomer shift is a linear function of the number of 5p "holes", V 9 5 ' 3 9 9 · On this basis, it has been concluded^ -397,399 that the bonds in I 2 , IC1, IBr, ICN, I2C16, IC12 ~ and IC14 - are derived almost exclusively from 5/7-electrons of the iodine with negligible ^-hybridization. By contrast, where sphybridization is involved, e.g. in I 0 4 ~ and CHwl4_w. (« = 0-3), ^-electron density is removed from the neighbourhood of the iodine nucleus into the bonding orbitals, and, as indicated in Fig. 22, the isomer shift relative to I _ is negative. In this case the isomer shift is empiri­ cally expressed by an equation of the type δ = -ah8 + bhp+c

where hs is the number of 5s "holes" and a, b and c are constants, which must be determined experimentally395,399,400.

8. Electronic and magnetic properties298'305'329,40i-405. Knowledge about the differ­ ences between the ground and excited electronic states can be gained from the electronic spectra of halide species; most widely and profitably studied have been spectroscopic transitions involving internal rearrangements of the dn electrons of transition-metal derivatives. For details of the ground and very low-lying excited states of transition-metal halides, such studies have been augmented by measurements of the magnetic properties. The energies of the Estates depend upon the balance of the following interdependent factors: (i) the interaction of the ^-electrons with the nucleus from which they are in­ completely screened; (ii) the electron-electron repulsion, variations of which for different states of a given dn configuration may be expressed in terms of the two Racah parameters B and C; (iii) the magnetic coupling of the spin and orbital momenta of the electrons, which may be defined in terms of the one-electron parameter ζ; and (iv) the field due to the ligands, which is usually referred to the ligand- or crystal-field parameter Dq. In the rearrangement of the dn electrons it is commonly assumed, as a first approximation, that the energy due 401 T. M. Dunn, Modem Coordination Chemistry (ed. J. Lewis and R. G. Wilkins), p. 229. Interscience, New York (1960). 402 H.-H. Schmidtke, Physical Methods in Advanced Inorganic Chemistry (ed. H. A. O. Hill and P. Day), p. 107. Interscience (1968). 403 A. B. P. Lever, Inorganic Electronic Spectroscopy, Elsevier (1968); J. Ferguson, Progress in Inorganic Chemistry, 12 (1970) 159. 4( 4 > B. N. Figgis and J. Lewis, Progress in Inorganic Chemistry, 6 (1964) 37; B. N. Figgis, Introduction to Ligand Fields, Interscience, New York (1966). 405 M. Gerloch and J. R. Miller, Progress in Inorganic Chemistry, 10 (1968) 1.

GENERAL PROPERTIES OF HALIDES

1277

to thefirstterm is unchanged, so that the study of d-d spectroscopic transitions bears witness to the effects of electron-electron repulsion, spin-orbit coupling406"408 and the ligand field.

Interelectron Repulsion: the Nephelauxetic Effect Comparison of the spectra of compounds formed by a given transition-metal ion with that of the gaseous metal ion shows that the Racah parameters are smaller for the com­ pounds than for the free ion, implying that coordination decreases the repulsions between ^/-electrons on the transition-metal ion. Decrease of the effective nuclear charge on the metal ^-electrons through covalent σ-bonding with the ligands causes an expansion of the d-orbitals, an effect which may be enhanced by metal-to-ligand ττ-bonding. Accordingly, the reduction in the Racah parameters, represented, for example, by the ratio ^compound/^free ion = ft c a n *>e usec * as an index to covalency in the metal-ligand bond. Ligands have been ordered on the basis of their increasing "nephelauxetic" effect (decreas­ ing ß) with respect to a given metal ion, e.g. F - < H 2 0 < N H 3 < H2NCH2CH2NH2 < SCN" < Cl- - CN" < Br~ < S2- ~ I"

which is also the order of increasing polarizability of the ligand atom.

Ligand-field Effects The ligand-field parameter Dq is well known to be a function of the metal ion, the ligand and the stereochemistry of the system. For a given metal and stereochemistry, the nature of the ligand causes Dq to increase in accordance with the familiar spectrochemical series, an abbreviated version of which follows: I" < Br~ < Cl" < S2- < F - < H 2 0 < SCN" < N H 3 < N 0 2 ~ < CH 3 ~ < CN"

Such an order cannot be rationalized on the basis of the simple point-charge model of crystal-field theory. In fact, the following factors all subscribe, in varying degrees, to the magnitude of Dq: interactions due to electrostatic perturbation, the metal-ligand σ-bond, metal-to-ligand ττ-bonding, ligand-to-metal ττ-bonding, and metal electron-ligand electron repulsions. For certain systems a scheme has been devised to resolve these various contri­ butions409: the following order of σ-bonding interaction with a given metal ion is thus found: NH 3 > H 2 0 > F " >C1" > B r " > I "

while the ττ-donor capacity follows the sequence I" >Br~ >C1" > F " > N H 3 406 B . N. Figgis, J. Lewis, F. E. Mabbs and G. A. Webb, / . Chem. Soc. (A) (1966) 1411. 407 j . Owen, Proc. Roy. Soc. A227 (1955) 183; Discuss. Faraday Soc. 19 (1955) 127. 408 j . M . Dunn, / . Chem. Soc. (1959) 623. 409 D . S. McClure, Advances in the Chemistry of the Coordination Compounds (ed. S. Kirschner), p. 498. Macmillan (1961).

1278

CHLORINE, BROMINE, IODINE AND ASTATINE: A. J. DOWNS AND C. J. ADAMS

9. Dipole Moments329»410»411. A neutral halide molecule in which the bonds to the halogen atoms are not symmetrically disposed inevitably possesses a dipole moment. For poly­ atomic halides, a bond moment can be defined for each bond, provided that the geometry of the molecule is known, though, as vector quantities, such moments are seldom strictly additive. In fact, the dipole moment of a unit M-X depends, not only upon the ionic character, but also upon additional factors: namely, (i) the overlap moment arising from the fact that the orbitals of two atoms of dissimilar size tend to overlap closer to the smaller atom; (ii) the hybridization moment due to the possible asymmetry of the atomic orbitals involved in the bond, as in PX 3 ; and (iii) the induced moment arising from the polarization of valence electrons. Only when allowance has been made for these other effects does the dipole moment provide a measure of the ionic character of diatomic halides which is substantially in accord with estimates based on nqr or electronegativity parameters387. More generally, however, the properties of M-X bonds in polyatomic halide species are not easily assessed in terms of dipole moments, though evidence of 7r-bonding has more than once been adduced in this way412. Dipole moments of halide molecules have been compiled elsewhere410»411. Nature of Bonding: a Summary The electronegativity of an atom being a measure of the energy of the valence electrons in relation to that of other atoms, the high electronegativity commonly associated with the halogens signifies the relatively low energy (in absolute terms) of the valence orbitals of these, compared with other, atoms. The binding of a diatomic halide is then determined by the relative energies of the valence orbitals of M and X, by the number and type of such orbitals, and by the number of valence electrons carried by M. In most circumstances only the outer ns- and «p-orbitals of the halogen appear to make significant bonding contribu­ tions, and, with respect to more electropositive elements, even the «^-orbital is predomin­ antly non-bonding in character; the participation of the vacant «d-orbitals of chlorine, bromine or iodine cannot be excluded, but the extent is now conceded by most theoretical studies to be small413. Relatively direct evidence that the "lone-pair" electrons of the halogen atom commonly occupy orbitals which are almost atomic in character comes from the recently described photoelectron spectra of halide molecules414; in simple monohalide derivatives of saturated hydrocarbons such orbitals have thus been shown to be virtually non-bonding, but in certain organic polyhalides and in numerous inorganic halides there is good reason to believe that these orbitals assume varying degrees of 7r-bonding character. For a typical element M in the same horizontal Period as X, the energy level diagram of MX, illustrated schematically in Fig. 23, varies from that characteristic of a highly polar alkali-metal halide molecule, with minimal perturbation of the appropriate atomic orbitals, to that characteristic of the exclusively covalent halogen molecule, the formation of which implies marked perturbation of the two separate sets of atomic orbitals. The bonding 4io B. Lakatos and J. Bohus, Ada Chim. Acad. Sei. Hung. 20 (1959) 115; B. Lakatos, J. Bohus, and G. Medgyesi, ibid. p. 1. 411 A. L. McClellan, Tables of Experimental Dipole Moments, Freeman, San Francisco (1963). 41 2See for example E. A. V. Ebsworth, Volatile Silicon Compounds, pp. 57-58,162. Pergamon, Oxford (1963); J. Lorberth and H. Nöth, Ber. 98 (1965) 969. 413 See for example B. M. D e b and C. A . Coulson, / . Chem. Soc. (A) (1971) 958. 414 D . W. Turner, C. Baker, A . D . Baker and C. R. Brundle, Molecular Photoelectron p. 214. Wiley-Interscience (1970).

Spectroscopy,

1279

GENERAL PROPERTIES OF HALIDES

orbitale of the MX molecule having a maximum capacity of 10 electrons, the highest dis­ sociation energies are exhibited by the Group III monohalides BF, A1C1, etc. The acceptor function of the halogen orbitals in the process of compound-formation is a characteristic of all halide species. Consequences of this electron-transfer include the M = alkali metal M σηρ

M

M

X

πηρ

X

σηρ ·:

*2

Atomic number (and electronegativity) of M

FIG. 23. Schematic energy level diagram for the diatomic halide MX of a typical element M belonging to the same period as X.

contraction and stabilization of outer orbitals on the atom partnered by the halogen; hence, vacant rf-orbitals on atoms such as silicon or phosphorus may be induced to play a significant part in σ- or ττ-interactions415. Even molecular-orbital accounts of electron-rich species like XeCl2 or IC14 -, which deny d-orbital involvement in the first instance, accentu­ ate the importance of charge-transfer to the halogen ligand, since it is here that the nonbonding molecular orbital of the three-centre-four-electron model (see Fig. 4) is largely localized. Theoretical accounts of the bonding in metal halides have been reviewed by Pearson and Mawby329. Total coordinate bond energies corresponding to the process MX»(g) -> M»+(g)+nX-(g) have been calculated on the basis of the following models: 415 D. P. Craig and E. A. Magnusson,/. Chem. Soc. (1956) 4895; D. P. Craig, Chemical Society Special Publication No. 12, p. 343. The Chemical Society, London (1958); K. A. R. Mitchell, Chem. Rev. 69 (1969) 157.

1280

CHLORINE, BROMINE, IODINE AND ASTATINE: A. J. DOWNS AND C. J. ADAMS

(i) an ionic model assuming spherical, non-polarizable ions; (ii) an ionic model incorporat­ ing polarizable ions; (iii) a localized molecular-orbital model in which a two-centre molecular orbital is constructed for each M-X bond; and (iv) a modified version of the semi-empirical Wolfsberg-Helmholz method, an LCAO-MO approach using the one-electron approxima­ tion. The hard-sphere model gives a good account of the alkali-metal halides, somewhat less satisfactory results for the alkaline-earth halides, and understandably poor results for compounds like AICI3 and TiCl4. The polarizable ion model works comparatively well for a wide range of metal halides, but gives poor results for some, e.g. BeCl2 and T1CI4. The energies so calculated are generally in fairly close correspondence with those derived from the localized molecular-orbital model, showing that both are approximate ways of evaluat­ ing the same effect, namely the distortion of the electron cloud of the anion in the field of the cation. The modified Wolfsberg-Helmholz method gives coordinate bond energies in good agreement with experiment for halides as widely different as NaCl and CCI4, but some of the assumptions of this treatment appear rather arbitrary and lacking in physical signi­ ficance. Though the complexity of the Hartree-Fock method at present restricts its applica­ tion to simple molecules containing atoms of relatively low atomic number, halides of Groups I-IV have been successfully treated by the extended Hückel method utilizing single-exponent, one-electron Slater orbitals416. Properties deduced in this way include the electronic structure, charge distribution, ionizatiön potential and bond angle. It has even been suggested that the ionic nature of the halides favours such calculations by permitting only a minimum of orbital overlap; relatively crude atomic orbitals are not, therefore, incompatible with a reasonable molecular-orbital scheme. 3.3. THE H Y D R O G E N H A L I D E S

Historical Background About AD 77, Pliny, in his Naturalis Historiae411', described the purification of gold by heating it with salt, "misy" (iron or copper sulphate) and "schistos" (clay). Although this mixture would give off fumes of hydrogen chloride, attention was focused on the effect of the treatment on the metal and not on the nature of the effluvia. There is reference in the "Alchimia Geberi"—possibly written as early as the thirteenth century—to the preparation of aqua regia418, while the preparation of hydrochloric acid itself was first reported in a fifteenth-century Italian manuscript419. Although there is no clear record, it is almost in­ conceivable that aqua regia or hydrochloric acid was not made earlier than this, since all the necessary ingredients were in the hands of early chemists. In Lavoisier's nomenclature, hydrochloric acid was designated "muriatic acid", a name still used in American industry420, but, following Davy's investigations on chlorine421, the name "hydrochloric acid" came into general usage. Reports of the preparation of hydrogen bromide and hydrogen iodide followed hard upon the identification of the corresponding elements. Thus, Balard showed 416 J. W. Hastie and J. L. Margrave, / . Phys. Chem. 73 (1969) 1105. ? Pliny, Naturalis Historiae, Book 33, chapter 25 (first century A.D.). 418 J. W. Mellor,/! Comprehensive Treatise on Inorganic and Theoretical Chemistry, Vol. II, Longmans, Green and Co., London (1922). 41

419 L . R e t i , Chymia,

420

10 (1965) 1 1 .

W. R. Kleckner and R. C. Sutter, Kirk-Othmer Encyclopedia of Chemical Technology, 2nd edn., Vol. 11, p. 307. Interscience (1966). 421 H . D a v y , Phil. Trans. 100 (1810) 2 3 1 ; Alembic

Club Reprints

(1894) 9.

THE HYDROGEN HALIDES

1281

that the passage of a mixture of hydrogen and bromine vapour through a red-hot tube containing iron turnings gives rise to hydrogen bromide422. Courtois first prepared hydro­ gen iodide without recognizing its nature, but, in his historic memoir on iodine, GayLussac showed that, although hydrogen and iodine do not react at ordinary temperatures, combination does occur in the vapour phase at elevated temperatures423. Preparation 1. Outline of methods345»418»424-428. The methods available for the preparation of the hydrogen halides, whether as gases or as aqueous solutions, are compared in Table 28. As indicated in the table, the methods fall into three general categories. The Direct Combination of the Elements The characteristics of this important and much studied reaction have been outlined in Section 2 (pp. 1223-4); more detailed accounts are to be found in references 345, 418 and 424-426. It has also been pointed out in preceding discussions that the relative ease of direct synthesis of the hydrogen halides varies with the bond energy of the HX molecule. Thus, the burning of hydrogen in chlorine, which takes place without catalytic agency, is a commercially important process for the production of hydrogen chloride. The less facile combination of hydrogen and bromine is most conveniently brought about in the presence of a catalyst (e.g. platinized asbestos, platinized silica gel or activated charcoal) at tempera­ tures between 200° and 400°C. Energetically least favoured is the formation from its ele­ ments of hydrogen iodide, though, with the aid of a platinum catalyst typically heated to at least 300°C, this probably affords the best method of preparing hydrogen iodide in other than small quantities. Reduction of the Parent Halogen by Agents other than Hydrogen Several of these reactions have been referred to incidentally in Section 2. From the pre­ parative standpoint the most important reducing agents are as follows. Red phosphorus and water, which furnishes a useful method of producing HBr or HI on the laboratory scale. Various hydrocarbons, the chlorination of which now provides much of the hydrogen chloride of commerce; the reaction of bromine or iodine with suitable unsaturated hydro­ carbons—notably tetrahydronaphthalene—represents a valuable laboratory route to the corresponding hydrogen halide. Sulphur dioxide: this is used to convert aqueous bromine to hydrobromic acid in one of the stages leading to the commercial production of elemental bromine (see pp. 1137-8). Hydrogen sulphide: the reaction of this with an aqueous suspension of iodine has been used to prepare aqueous hydriodic acid. 422 A. J. Balard, Ann. Chim. Phys. 32 (2) (1826) 337. 423 B . Courtois, Ann. Chim. 88 (1) (1813) 304, 311; J. L. Gay-Lussac, ibid. 91 (1) (1814) 5. 424 Gmelins Handbuch der Anorganischen Chemie, 8 Auflage: "Chlor", System-nummer 6, Verlag Chemie, Berlin (1927); "Brom", System-nummer 7, Verlag Chemie, Berlin (1931); "Iod", System-nummer 8, Verlag Chemie, Berlin (1933). 425 Gmelins Handbuch der Anorganischen Chemie, 8 Auflage, "Chlor", System-nummer 6, Teil B— Lieferung 1, Verlag Chemie, Weinheim/Bergstr. (1968). 426 Supplement to Mellows Comprehensive Treatise on Inorganic and Theoretical Chemistry, Supplement II, Part I, Longmans, Green and Co., London (1956). 427 G. Brauer, Handbook of Preparative Inorganic Chemistry, 2nd edn., Vol. 1, Academic Press (1963). 428 inorganic Syntheses, Vols. 1, 3 and 7, McGraw-Hill (1939-63).

J2 go

(e) with hydrocarbons

HC1 produced as a by-product of the chlorination of hydro­ carbons, e.g. CH 4 , CgHö. Com­ mercially a major source of H O

Occurs in the gas phase or in solution; S2CI2 also tends to form Reaction violent and inconvenient

(c) with hydrogen sulphide

(d) with red phosphorus and water

15-90°C, in aqueous solution

Provides a well-known laboratory method of preparing HBr Reaction of Br2 with benzene, toluene, naphthalene or petro­ leum (in the presence of a catalyst) gives HBr. Rapid re­ action with tetrahydronaphthalene (tetralin) at 20°C is a useful method of preparing HBr in the laboratory

One of the stages in the commercial production of bromine (in aque­ ous solution) Reaction tends to give S2Br2 as well as HBr

Has been used to prepare hydnodic acid, H2S being passed into an aqueous suspension of I2 Provides a convenient laboratory method of preparation Few hydrocarbons react readily but the reaction of I2 with boiling tetrahydronaphthalene has been used to prepare HI in the laboratory

700-800°C, over charcoal. HI solutions also produced by heating an aqueous suspension of I2 with activated char­ coal Reaction occurs but has been little exploited

In ultraviolet only Used as a method of synthesis, usually on a relatively small scale

200°C Principal method of synthesis on the large scale

20°C Important method of synthesis on the large scale

500°C, over charcoal

300°C* 600°C*

150°C 250°C

20°C

500°C, over charcoal

500°C*

HI

250°C 500°C

HBr

100°C 200°C

HC1

(b) with sulphur dioxide

B. Reduction of the parent halogen by agents other than hydrogen: (a) with water vapour, rapid reac­ tion at

A. Direct synthesis from the elements'. (a) without catalyst: slow combination at rapid combination at (b) with catalyst: slow combination at rapid combination at (c) under the influence of visible light: rapid combination at Comments

Method of preparation

TABLE 28. PREPARATION OF THE HYDROGEN HAUDES*"'

00

E.g. PBr3 (see above), AlBr3. The hydrolysis of AlBr3 has been used to prepare HBr of high purity

E.g. PCI3, TiCl4. Hydrolysis of MgCl2, CaCl2 and NaCl has been utilized commercially. Hy­ drolysis of organic acid chloride useful for the preparation of DC1 Under suitable conditions, H2 or simple hydrocarbons react with many metal chlorides to give HC1; of little practical use as methods of preparation Heavy-metal bromides reduced to anhydrous HBr by either H2 or a hydrocarbon at elevated tem­ peratures. H2S may be used, as in CdBr 2 +H 2 S -> CdS + 2HBr; of negligible practical signifi-

H2SO4 causes some decomposition through oxidation of the HBr; H3PO4 gives satisfactory results H2SO4, CaBr2 or P 2 0 5 may be used as dehydrating agents

"Salt-cake" process using H2SO4 and NaCl still of considerable industrial importance A useful method of preparing gaseous HC1 in the laboratory

H 2 S may be used, as in 2CuI+H 2 S-> CU2S+2HI; of negligible practical significance

E.g. PI3, Sil 4

P2O5 may be used as a dehydrating agent

H2SO4 causes extensive oxidation of HI; H3PO4 can be used satisfactorily

* Extensive thermal decomposition of HI occurs at temperatures above 300°C, so that an equilibrium mixture of HI, H2 and I2 is obtained. a Gmelins Handbuch der Anorganischen Chemie, 8 Auflage: System-nummer 6, "Chlor" (1927); "Chlor", Teil B, Lieferung 1 (1968); System-nummer 7, 'Brom" (1931); System-nummer 8, "Iod" (1933). b J. W. Mellor, A Comprehensive Treatise on Inorganic and Theoretical Chemistry, Vol. II, Longmans, Green and Co., London (1922). c Supplement to Mellor's Comprehensive Treatise on Inorganic and Theoretical Chemistry, Supplement II, Part I, Longmans, Green and Co., London (1956). d Z. E. Jolles (ed.), Bromine and its Compounds, Benn, London (1966). e G. Brauer, Handbook of Preparative Inorganic Chemistry, 2nd edn., Vol. 1, Academic Press (1963). f Inorganic Syntheses, Vols. 1, 3 and 7, McGraw-Hill (1939-63).

(d)by reaction with H2, hydrocarbons or H2S

(b) action of concentrated H2SO4 or other dehydrating agents on a concentrated aqueous solu­ tion of HX (c) by hydrolysis

C. From inorganic or organic halides (a) action of concentrated H2SO4 or H3PO4 on metal halide

1284

CHLORINE, BROMINE, IODINE AND ASTATINE: A. J. DOWNS AND C. J. ADAMS

From Inorganic Halides The displacement of the hydrogen halide from a metal halide by the action of a less volatile protic acid forms the basis of the classical method of producing hydrogen chloride on both the laboratory and the commercial scale. Thus, the reactions NaCl+H 2 S0 4 -> NaHS04+HCl

and

NaHS04+NaCl -» Na 2 S0 4 +HCl

make up the so-called "salt-cake" process, which, from its inception as part of the Leblanc process, has been used for over a century in the manufacture of hydrogen chloride. By contrast, the oxidizing action of hot, concentrated sulphuric acid militates against the pre­ parative success of analogous reactions involving a metal bromide or iodide, since the hydro­ gen halide tends to be oxidized to the parent halogen. Nevertheless, the use of syrupy phosphoric acid in place of sulphuric acid obviates this difficulty. The dehydration of concentrated aqueous solutions of the hydrohalic acids, either by chemical agents like sulphuric acid or phosphorus pentoxide or by thermal stripping, also serves as a convenient source of the gaseous hydrogen halide. The hydrolysis of certain halides has likewise been turned to advantage, as in the preparation of pure hydrogen bromide by the hydrolysis of aluminium bromide, or in the production of hydrogen chloride as a by-product resulting from the hydrolysis of magnesium chloride (in the Dow process for extracting magnesium from sea water). The hydrolysis of organic acid chlorides, e.g. QHsCOCl, with deuterium oxide affords one of the most convenient methods of preparing deuterium chloride427. 2. Laboratory methods345»424 ~430. Although the aqueous hydrohalic acids are readily available from commercial sources, pure samples of a gaseous hydrogen halide are less easily obtained in this way. Of the various methods which have been suggested for the pro­ duction of the gaseous compounds, the following are probably the most effective: Hydrogen chloride: (i) the action of concentrated or pure sulphuric acid on sodium chloride; (ii) the action of concentrated sulphuric acid on concentrated (e.g. constantboiling) hydrochloric acid. Hydrogen bromide: (i) direct union of the gaseous elements in the presence of a catalyst; (ii) the action of bromine on a mixure of red phosphorus and water represented mainly by the equation 2P+6H20+3Br2 -> 6HBr+2H3P03 (iii) the reaction of bromine with tetrahydronaphthalene at room temperature

(fV >

+.4ΒΓ2-

™f

T

T+4HB,

though this involves the loss of at least half of the bromine taken; (iv) dehydration of a con­ centrated aqueous solution of hydrobromic acid, e.g. with phosphorus pentoxide, concen­ trated sulphuric acid431 or calcium bromide. 429 R. E. Dodd and P. L. Robinson, Experimental Inorganic Chemistry, pp. 197-200. Elsevier (1954). 430 R. H. Herber (ed.), Inorganic Isotopic Syntheses, Benjamin, New York (1962). « I A. D. B. Sloan, Chem. and Ind. (1964) 574.

THE HYDROGEN HALIDES

1285

Hydrogen iodide: (i) direct union of the gaseous elements in the presence of a catalyst; (ii) the action of iodine on a mixture of red phosphorus and water; (iii) the iodination of boiling tetrahydronaphthalene; (iv) dehydration of a concentrated aqueous solution of hydriodic acid. Practical details are given in references 424 and 426-431. Gaseous hydrogen chloride is freed from sulphur dioxide by scrubbing with sulphuric acid, which also serves as a drying agent. Hydrogen bromide is stripped of elemental bromine by passage over moist red phosphorus or through tetrahydronaphthalene; phosphorus acids and bromides are ab­ sorbed in a small quantity of water, while the best drying agents appear to be anhydrous calcium bromide and activated alumina. Iodine is removed from hydrogen iodide by con­ densation or by reaction with a metal iodide, either as a solid or as a saturated aqueous solution; anhydrous calcium iodide and phosphorus pentoxide are recommended drying agents. Samples of hydrogen halide of high purity, suitable, for example, for liquid-phase studies of their solvent properties432, are obtained by repeated fractionation in vacuo. A guide to the purity of the material is provided by its vapour pressure, though the most sensitive test is said to be afforded by the specific conductivity of the liquid432. Of the three compounds, hydrogen iodide presents the most taxing problems of manipulation since it is decomposed by light, by many organic materials, e.g. tap-grease, and also by extensive glass surfaces, particularly in the presence of traces of water. An aqueous solution of a hydrogen halide free from impurity is readily obtained by absorption of the purified gas in water. For exceptional purposes, very pure hydrochloric acid has been prepared by isothermal distillation of a concentrated solution of the acid in a desiccator containing a sample of scrupulously purified water425; an alternative pro­ cedure involves boiling the acid with small amounts of potassium permanganate to remove traces of bromine and iodine, followed by distillation via SL quartz condenser427. Constantboiling hydrobromic acid is conveniently prepared from potassium bromide, water and concentrated sulphuric acid, while constant-boiling hydriodic acid can be made by the reduction of iodine in aqueous suspension with either hydrogen sulphide or hypophosphorous acid428. Concentrated solutions of hydrobromic acid are somewhat susceptible to oxidation, and, unless protective measures are taken, the formation of bromine and polybromide ions causes pronounced darkening over a period of time. The tendency to suffer oxidation is even more pronounced in the case of hydriodic acid, the decomposition of which is favoured by the action of light and of impurities; storage in a well-stoppered, dark-glass bottle is recommended to preserve the acid in a satisfactory condition for ex­ tended periods. The addition of small amounts of hypophosphorous acid or of red phos­ phorus inhibits the oxidation of hydriodic acid. Ion-exchange methods have proved effective in the purification of hydrobromic and hydriodic acids433. Methods used to prepare isotopic variants of the hydrogen halides include345»425 ~430 (i) the direct interaction of deuterium or tritium with the halogen; (ii) hydrolysis with deuter­ ium oxide or tritium-enriched water of a suitable halide, examples being PX3 (X = Cl, Br or I, either as the pure compound or formed in situ by the reaction of the elements), PCI5, S1CI4, SOCl2, MgCl2, AICI3 and C 6 H 5 COCl; (iii) the reaction of dry sodium chloride with deuterosulphuric acid; (iv) the reduction with deuterium or tritium of a metal halide, 432 M . E. Peach and T. C. Waddington, Non-aqueous Solvent Systems (ed. T. C. Waddington), p. 83. Academic Press (1965). 433 H. Irving and P. D. Wilson, Chem. and Ind. (1964) 653.

1286

CHLORINE, BROMINE, IODINE AND ASTATINE: A. J. DOWNS AND C. J. ADAMS

e.g. C0CI2 and AgCl; and (v) exchange reactions, as between concentrated aqueous HC1 and tritium-enriched water or sulphuric acid, or between tritium and HC1 at elevated tem­ peratures. Labelling by introduction of a radioactive halogen isotope has also been accom­ plished by one or other of these methods, e.g. D 2 S0 4 +Na36Cl -> NaDS04+D36C1428

while thermal diffusion of the gaseous hydrogen halide affords the most effective means of separating the naturally occurring species H35C1 and H37C1 and likewise H79Br and H81Br (seep. 1150). Manufacture420,434,435 Hydrogen chloride and hydrochloric acid are commodities of much commercial impor­ tance, world production being in excess of 3,000,000 metric tons of hydrogen chloride per annum. By contrast, hydrogen bromide and iodide are of small commercial significance, whether in the gaseous or in the aqueous states. The commercial production of hydrogen chloride awaited the Leblanc process for the manufacture of sodium carbonate, in which the gas was a co-product with salt-cake (sodium sulphate) of the first step, the reaction between salt and sulphuric acid. For a time, the gas was merely vented to the atmosphere, but legis­ lation was enacted prohibiting its indiscriminate discharge, and thus necessitating its recovery. Although the Solvay process has supplanted the Leblanc process, the reaction between salt and sulphuric acid is still exploited because of the industrial demand for saltcake (e.g. by the paper and glass industries) and hydrochloric acid. At the present time, four major processes are used commercially to produce hydrogen chloride and hydro­ chloric acid: (i) the salt-sulphuric acid process, (ii) the Hargreaves process, (iii) the direct synthetic method, and (iv) organic chlorination processes, which yield hydrogen chloride as a by-product. Usually two or more of these processes are operated in every major indus­ trialized country of the world, depending upon the availability of raw materials and the size and diversity of the chemical industry of that country. Salt-sulphuric Acid Process The endothermic reactions and

NaCl + H2SO4 -> NaHS0 4 +HCl NaCl+NaHS0 4 -> N a 2 S 0 4 + H C l

take place at temperatures in the order of 150 and 540-600°C, respectively. In practice, temperatures are usually kept below 650°C to prevent the salt-cake from fusing. The reactions are carried out in various types of equipment: for example, a cast-iron retort; the so-called Mannheim furnace, a mechanical furnace consisting of a stationary circular muffle in the form of a basal concave pan and a domed cover, separated by a cylindrical mantle; the Laury-type furnace, which has a mobile oil-fired combustion chamber and, for the reaction vessel, a horizontal two-chambered rotating cylinder; or the Cannon fluid-bed 434 V. A. Stenger and G. J. Atchison, Kirk-Othmer Encyclopedia of Chemical Technology, 2nd edn., Vol. 3, p. 767. Interscience (1964). 4 35 A . W. Hart, M. G. Gergel and J. Clarke, Kirk-Othmer Vol. 11, p. 857. Interscience (1966).

Encyclopedia

of Chemical Technology, 2nd edn.,

THE HYDROGEN HALIDES

1287

reactor, wherein sulphuric acid vapours are injected with the gases from a gas-fired com­ bustion chamber into a fluid bed of salt. Many references in the patent literature relate to improvements in the efficiency of the process, by modifications of temperature, relative quantities of reactants, and design of reaction chambers. In some cases, potassium or calcium chloride has been employed in place of sodium chloride. Hargreaves Process Here the reactants are salt, sulphur dioxide, air and water vapour; the products are the same as in the salt-sulphuric acid process. The essential reaction 4NaCl+2S0 2 + 0 2 + 2 H 2 0 - * 2Na 2 S0 4 +4HCl

is exothermic, and sufficient heat is evolved to maintain the process once the reactants have been brought to temperatures in the order of 430-540°C. In practical terms, a mixture of sulphur dioxide, steam and air is passed through stacks of salt briquets lying on perforated trays within a vertical chamber. Introduced in England in the last half of the nineteenth century, the Hargreaves process promised for a time to become a major source of salt-cake and hydrochloric acid, but various factors have led to its decline in more recent times. Direct Synthetic Process Hydrogen chloride in high concentration and of high purity is synthesized by the com­ bustion of a controlled mixture of hydrogen and chlorine. The chlorine is derived, either wet or dry, from the electrolysis of brine; the hydrogen may come likewise from the electro­ lysis of brine, or from other sources, e.g. the hydrocarbon-steam reaction. The equipment varies in detail with the qualities of the raw materials and of the product desired, but the essential parts comprise a burner and combustion chamber, necessary control and safety devices, and facilities for processing the hydrogen chloride. The burner consists of a nozzle that injects the reactants into a vertical or horizontal combustion chamber; materials favoured for construction are silica, brick-lined steel, water-jacketed steel or water-cooled graphite. The reaction is initiated by igniting the hydrogen in a stream of air with a retract­ able torch burning either hydrogen or coal-gas; chlorine is then introduced into the burner to establish a hydrogen-chlorine flame. To ensure the formation of chlorine-free hydrogen chloride, a slight excess (typically 2-5%) of hydrogen is employed. By-product Hydrogen Chloride The chlorination of many organic materials produces hydrogen chloride as a by-product. In recent years the scale of chlorination processes of this sort has reached such proportions that in the United States, for example, more than 75% of all the hydrochloric acid of commerce is so derived. The chlorination of methane or benzene is typical of the reac­ tions which generate hydrogen chloride as a by-product; the gas is also produced by pyrolysis of certain chlorinated organic compounds, as in the cracking of 1,2-dichloroethane or 1,1,2,2-tetrachloroethane: C1CH2CH2C1 C12CHCHC12

► CH 2 = CHC1 + HC1 catalyst

► CCJ2 = CHC1+HC1

The hydrogen chloride evolved from these reactions is liable to be contaminated with chlorine, air, organic chloro-compounds, excess reactants and moisture, depending upon

1288

CHLORINE, BROMINE, IODINE AND ASTATINE: A. J. DOWNS AND C. J. ADAMS

the individual process. Accordingly, extensive treatment is commonly necessary to strip the hydrogen chloride of its impurities. As noted previously, hydrochloric acid is also a by-product of the Dow process for extracting magnesium from sea water. Gas Treatment After leaving the generating plant, the hydrogen chloride is treated in several steps· These may include three or more of the following: (i) Removal of suspended solids in settling chambers, centrifugal traps or cyclone-type separators, (iij Cooling by heat-exchange methods, (iii) Absorption in water, usually achieved by counter-current flow of gas and liquid, (iv) Desorption of gaseous hydrogen chloride from the concentrated aqueous acid, usually brought about by thermal stripping at 75-130°C, though chemical agents, such as concentrated sulphuric acid or calcium chloride, have also been employed, (v) Purification, Methods used to purify by-product hydrogen chloride typically involve scrubbing of the gas with one or more non-aqueous solvents; alternatively the hydrogen chloride is removed from the gaseous products and purified by an absorption-desorption process using the aqueous acid as solvent. Activated charcoal has also been used to remove organic impur­ ities. (vi) Liquefaction by compression and cooling of the anhydrous gas. The exact treat­ ment depends on the composition and temperature of the raw gas and on the composition and nature of the end-product. Materials of Construction The choice of materials for the fabrication of equipment allowing the manipulation, storage and transport of hydrochloric acid is conditioned by the highly corrosive nature of the solution. Even with relatively dilute solutions, the corrosion problem is severe, and the situation is further complicated by the presence of contaminating materials, originating in the preparative process, which accelerate corrosion rates and shorten the lifetime of equipment. In the production of reagent-grade acid, the choice of materials is practically limited to tantalum, glass and impervious graphite. Where contamination can be tolerated, an economic balance must be struck between the life of the material and its cost. The follow­ ing materials have found extensive use: rubber-lined carbon steel, high-silicon iron alloys (e.g. "Durichlor"), nickel alloys (e.g. "Chlorimet" and "Hastelloy"), copper alloys, tan­ talum, titanium, acid-proof brick, chemical stoneware or porcelain, glass, certain rubbers, plastics (e.g. polyvinyl chloride, polyethylene, polystyrene, polytetrafluoroethylene and glass-reinforced polyesters), baked carbon, graphite or impregnated carbon and graphite. The limitations on the use of metals do not apply to the anhydrous gas; only in the presence of moisture do the corrosion problems become severe. In fact, no ignition point has been found with anhydrous hydrogen chloride in contact with steel even at 760°C, and the gas or liquid is usually stored in steel vessels. Uses Hydrochloric acid is widely used industrially in such diverse fields as the pickling of metals for scale removal, the reactivation of bone charcoal and carbon in sugar-refining operations, the production of glucose and corn sugar from starch, the production of chloroprene and vinyl chloride—important intermediates in the manufacture of synthetic rubber— and the activation of petroleum wells. Other applications are found in the production of alumina for the manufacture of aluminium; in the Dow process for isolating magnesium from sea water; in many extractive metallurgical processes for treating high-grade ores, among which are those yielding radium, vanadium, tungsten, tantalum, manganese and

THE HYDROGEN HALIDES

1289

germanium; in the production of numerous organic chloro-derivatives and metal chlorides; in the production of phosphoric acid from phosphate rock; as a catalyst in organic syn­ theses; and as an analytical reagent. In 1963 it has been estimated that, of the hydrochloric acid consumed, nearly 50% went to the synthesis of organic chemicals, 17% to the produc­ tion of metals, 18% to the activation of petroleum wells, 7% to metal and industrial cleaning operations, 5% to the production of inorganic chemicals, and 4% to food-processing. However, oversupply in recent years, combined with the increasing pressure to reduce pollution, has stimulated the search for new industrial outlets for the acid. Two develop­ ments which may assume importance in this respect are the production of chlorine from hydrochloric acid, either by electrolysis or by chemical oxidation (see Section 2), and "oxyhydrochlorination" processes, whereby hydrogen chloride, air or oxygen, and a hydro­ carbon give rise to chlorinated organic products and water, e.g. C 6 H 6 + H C H - i 0 2 -> C6H5C1 + H 2 0

Hydrobromic and Hydriodic Acids Hydrogen bromide gas is prepared commercially by burning a mixture of hydrogen and bromine vapour. The gas is passed through hot activated charcoal to remove free bromine and is either liquefied or absorbed in water. The major use of aqueous hydro­ bromic acid is in the manufacture of inorganic bromides and of certain organic bromoderivatives. Neither hydrogen iodide nor hydriodic acid appears to be produced on the large scale or to find significant use outside the laboratory. Physical Properties289,345,424 -426,432,436

1. General physical characteristics. Quantitative physical properties of the anhydrous hydrogen halides are summarized in Table 29, which illustrates the relatively uniform pattern of behaviour in the series HC1, HBr, HI; in many respects, however, this uniformity is not shared by hydrogen fluoride (see Chapter 25). Thus, unlike hydrogen fluoride, which boils at 19-5°C, hydrogen chloride, bromide and iodide all boil well below room temperature, the order of the boiling points being HF > HC1 < HBr < HI. The strongly hydrogenbonded fluoride apart, the relative volatility of the compounds reflects the strengthening of the van der Waals' forces which accompanies the increasing number of electrons, and hence the polarizability, in the series Cl, Br, I. Since the dipole moment of the molecules decreases progressively from HF to HI, dipole-dipole interactions presumably play a less important part than do dispersion or dipole-induced-dipole interactions in the intermolecular binding of the heavier hydrogen halides. On the evidence of the boiling points, Trouton constants and spectroscopic and other properties, hydrogen chloride, bromide and iodide differ from the fluoride in showing no signs of appreciable association in the liquid or gaseous states. The HX molecules persist throughout the solid, liquid and gaseous phases, though, as discussed below, hydrogen-bonding probably makes a significant contribution to the inter­ molecular binding of the solid. The pattern of behaviour thus exhibited by the hydrogen halides finds a close parallel in the properties of the hydrides H 2 0 , H 2 S, H2Se and H 2 Te and, to a lesser degree, in those of the corresponding Group V hydrides, each series being distinguished by the anomalous properties occasioned by the associative tendencies of the first member. 436 T. C. Waddington, MTP International Review of Science: Inorganic Chemistry Series One, Vol. 3 (ed. V. Gutmann), p. 85, Butterworths and University Park Press (1972).

Thermodynamic properties AHf° at 298°K (kcal mol"*) AGf° at 298°K (kcal mol"i) 5° at 298°K (cal deg"i mol~i)

^298°K

Force constant, fce(mdyne/Ä) Dissociation energy: D(


Anharmonic vibrational constant,

Ground state properties Internuclear distance, r«(A) Dipole moment, D Molecular quadrupole moment ( x 1026esucm2) Vibrational frequency, a^cm" 1 )

Gaseous molecules Electronic ground state configuration a1-** Observed transitions in the ultraviolet,15 vo,o(cm~i) C+-X 1Π«-1Σ + (Χ) B^X A«-X

Molecular weight

Property

5-161°

-22-063* -22-778* 44-643*

kcal 102-24* 103-16*

eV 4-433* 4-473*

H35Q, 52-5707;° H37C1, 52-5166° D35C1, 26-92;°.d D37Q, 26-99°«d

kcal 86-64* 87-54*

- 8-70* -12-77* 47-463*

4-114«

eV 3-757* 3-796*

H79ßr, 45-2316;* HSißr, 45-2175* D7°Br, 22-73; h D^ißr, 22-72*

5-8f 5-5f H35C1, 2990-7027;° H37Q, 2988-4799° H7QBr,2649-3876;*H8iBr, 2648-9752^ D79ßr, 1885-33 ; h D8iBr, 1884-75h D35C1, 2144-77;d D37Q, 2140-20c

l-4145 d 0-828e

70,542 67,084 Continuous absorption starting at ~ 35,000 cm" 1 with maximum at -55,000 cm-i(€ m a x 525)

77,540 75,134 Continuous absorption starting at 44,000 cm" 1 with maximum at -65,000 c m - i ( € m a x 890) l-2746°.d 1-07*

σ27τ3σ* ·*— σ 2 π 4

σ2π3σ*·<— σ 2 π 4

a

1Σ +

80-912

HBr

1Σ +

36-461»

HC1

TABLE 29. PHYSICAL PROPERTIES OF THE HYDROGEN HALIDES

kcal 70-43* 71-32* + 6-30* +0-38* 49-350*

3-1381

Hi 2 7I, 38-9810* Di 2 7I, 19-8731

—

eV 3 054* 3 093*

H ^ I , 2308-901* Di 2 7I, 1639-655*

l-6090 d 0-448e

62,320 56,750 Continuous absorption starting at 27,500 c m - 1 with maximum at -48,000 cm-i(«r max 174)

σ27τ3σ* <— σ27τ4

ΐΣ +

127-9125 a

HI

Liquid

Critical temperature Critical pressure (atm) Critical density (g cm" 3 ) Vapour pressure, p (mm Hg) Solid

Properties of the condensed phases Boiling pointi Melting pointi Liquid range at normal pressures (°C) Heat of vaporization at boiling point, AjFfvaP(kcalmol-i) Entropy of vaporization at boiling point (Trouton constant) (cal deg"i mol~i) Heat of fusion (kcal mol"i)

lonization potentials :k ΗΧ + ( 2 Π 3 / 2 )«-ΗΧ(ΐΣ + ) ΗΧ + (2Πι / 2 )«-ΗΧ(ΐΣ + ) ΗΧ+(2Σ+)^-ΗΧ(ΐΣ+)

°K 324-7c 8207 c 0-424c

°C 51-54c

—

84 m

°C 89.9m

°C -66-73 -86-88

eV 11-67 1200 15-29

°K 424.1m

°K 237-80 222-36

kcal 239-4 254-8 319-4

—

82*

19-861 0-6861

4-7241

15-441

°C 150-9m

°C -35-36 -50-80

eV 10-38 11 05 13-85

log/>= - ( 1 2 5 0 / Γ ) - 0-2766 log T l o g / ? = - ( 1 4 0 6 / D -- 0-377 log T - 0004084Γ - 0003167Γ + 10-493h + 10-488n log/7 = -(1338/Γ) -■ 4-672 log T l o g p = - ( 1 6 3 6 / D -- 7111 logT + 20-179 n + 0002293Γ + 26119 h

°K 363-l m

20-391 0-5751

20-51 0-4761

20-151 4-2101

°K 206-43 * 186-28

kcal 269-1 276-7 352-6

3-8601

°C - 85-05 -114-22

°K 188-11 158-94 29-171

eV 12-74 12-82 16-23

kcal 293-8 295-7 374-3

logp = - ( 1 1 1 4 / Γ ) - 1-285 log T - 00009467Γ + 12005 (130-160°K) h log/> = -(905-53/Γ) + 1-75 log T - 0005077Γ + 4-65739h

1

a-phase (stable below 98-38 °K)°

HC1

Thermal properties Heat capacity, Cp(cal mol ~1 deg " i) Solid Liquid Gas

0-340-11-69 (10-76-147 ·4°Κ)* 14-41-14-44 (163-0-173-4°K)c 6-959-9-388 (100-6000°K)*

|

Each hydrogen halide exists in polymor­ Face-centred orthorhombic lattice, space group Bb2\m. 4 molecules phic crystalline forms, of which that per unit cell. Unit cell dimen­ stable at low temperatures has been most sions for DC1 at 92-4°K (Ä) fully characterized. The incomplete and a b c not always concordant evidence of 5-410 various physical measurements on this 1 5082 5-826 form points to the presence of zigzag
Crystal structure

Property HBr

Low-temperature form (stable below ca. 120°K)8

HI

Hsißr, 447-76 DSißr, 443-39 T*iBr, 441-26 Ramanu4»u6-U7

1-831-12-32 (15-72-182-09°K)V 14-20-14-31 (189-93-205-ll°K) v 6-959-9-422 (100-6000°K)X

H79ßr, 535-44 D79ßr, 530-58 T79ßr, 528-75 Infrared ;u* ~ u * u 7

2-745-12-04 (17-11-218-00°K)W i4.34-i4.i5 (227-05-236-15°K)w 6-959-9-511 (100-6000°K)X

H127I, 183107 D127I, 1823-38 T127I, 1822-61 Infrared"*

Face-centred orthorhombic lattice. Face-centred tetragonal lattice. 4 molecules per unit cell. Unit 4 molecules per unit cell. Unit cell dimensions at 4-2°K (A) cell dimensions for DBr (Ä) a c\a Temp. a b c 5-987 1082 84°K 5-44 5-614 6-120* 107°K 5-559 5-649 6-120f High-temperature form (stable above ca. 120°K)8 d ( D - B r ) ~ 1-38 ± 0 0 3 Ä Z ( B r - - B r - · Br) = 91° 41'* or Face-centred cubic lattice. 4 mole­ 90° 56'f cules per unit cell. Lattice * Phase isomorphous with a-phase of DC1, space group Bb2\m. parameterat 1 3 0 ° Κ , Λ 0 = 6-246Ä Ferroelectric.13 t Disordered phase stable between 86° and 118°K, space group Bbcm High-temperature form (stable above ca. 1 1 0 ° ^ Face-centred cubic lattice 4 molecules per unit cell Lattice parameter ao = 5-76-5-78 A

Low-temperature form(s) (stable below ca. 110°Κ)Γ

TABLE 29 (cont.)

Gas at N.T.P. Compressibility ( x 106 per bar) Solid Liquid Viscosity, i?(centipoise) Liquid Gas Surface tension of liquid, y(dyne c m - 1 )

Mechanical properties Density, d(g cm" 3 ) Solid Liquid

Coefficient of cubical expansion ( x 104 per °C) Solid Liquid Gas

Gas

Liquid

Thennal conductivity (cal s e c - 1 c m - 2 deg" 1 cm" 1 ) Solid

0-874-0-816 (186-8-199-4°K) ii 0-01835-0-02365 (291-9-373-4°Κ)* 30-191-25-399 (182-204°K)8*

0-566-0-058 (160·8-318·2°Κ)" 001313-002530 (273-2-523-2°^ 26-912-22-409 (168·5-192·6°Κ)«

—

ee

36-5-7-4 (0-20 kbar, 150°K)

ee

h

2-96-2-55 (0-150°K) h ' ee d= 2-757 - 1-238 X 1 0 - 3 Γ 7 0 1 x 10-6J2 ( Γ = 193-333°K) ff 3-6445 x 10-3 eg

21-19 (206-1-227-2°K)

5.4CC

1-24-4-29 X 10-5 (200-600°K;/7 = 130-250 mm Hg) aa

0-625-1-321 X 10" 3 (780-142-2°K) y

3 6 0 - 7 0 (0-20 kbar, 130°K) 380-53 (100-2000 bar, 293°K)C

1-6392 x 10-3 c

1-54-1-47 (0-107°K) h 1-267-0-630 (160-01-319-35°K)C

6-2° 27-8-24-0 (159-188°K)C 37-69-36-61 (273-373 °K, p = 0-1 atm)c

8-71-1-66 X 10-4 (199-2-337-2°K; pressure 30-200 bar)z 2-18-6-73 X 10-5 (200-600°K;/> = 210-260 mm Hg) aa - bb

—

1-418-1-298 (223·2-236·4°Κ)" 001731-003228 (273-2-524-2°K) u 29-06-26-96 (225-3-236-5°K) hh

— —

5-7888 x 10-3 hh

3-67-3-05 (4-2-222°K) 8 ' hh 2-860-2-7862 (223-3-240-4°K) hb

5 (4-2-180°K) s 14-4 (237-8°K)dd

—

—

Molar refraction (cm* mol" 1 ), D-line

^

6-666*

nD = 1-254-1-252 (283-290°K)« λ = 6562-9-2302-83 Ä n = 1-0004434-1 000523525c.h

3-5 x 10-9 (l88°K)c»h Temp.coefficient = - 6 - 5 x 10"H ( o h m - 1 c m - 1 deg" 1 )

Liquid

Refractive index, n Liquid Gas at N.T.P.

0-13-1-3 X 10-9 (87-1-99-1 °K)C <10-ΐ0(99·1-159·0°Κ)°

a-phase: 2-6-4-3.° Slight dielectric absorption. ß-phase: Static dielectric constant €0 = 24-7-15-2 (98-4-158-9°K)c Dielectric absorption gives com­ plex dielectric constant. 14-30-4-6 (158-301 °K)C Some dielectric absorption e = 1 007452-1 001182 (201-4588-8°K,/> = latm) c ' h c = 1 00526-1 -24474 (p = 0-993-39-936 atm, T = 298°K)C = 1-00476-1-18271 (/? = € 0-995-40-210 atm, Γ=323°Κ)°

HC1

Solid

Specific conductivity, k (ohm~i cm - 1 )

Gas

Liquid

Electrical, optical and magnetic properties Dielectric constant, c Solid

Property

HBr

9-161™

n D = 1-325 (283 °K)** λ = 6707-85-4799-9 A n = 1 00060752-1 00062160"

1-4 X 10-9 (189°K) mm

—

e = 1 004545-1 000943 (217-5599-1 °K,p = 1 atm)*

Static dielectric constant e0 = 200-8-23 (70-187°K; 20-1-3 x 108 Hz).kk»n Dielectric absorp­ tion gives complex dielectric constants; behaviour very sensi­ tive to phase transitions. 7-33-3-82 (188-298°K) n

TABLE 29 (cont.)

Conductivity x 10io v= v — 60 kc/s 300 c/s 6-8-1-7 <0-l <1 <01 3-4-^-2 <01-0-6 4-2-13 10-3-4 (ref. hh)

13-73™

nD= 1-466 (289-7°K)** λ = 6707-5-4799-9 Ä n = 1 00091087-1 •00093900**

Temp. range, °K 85-7-118-5 126-3-200-2 209-6-222-4 222-6-236-3

€ = 1-002687-1-0009706 (244-5612-2°K,/> = latm)**

Static dielectric constant e0 = 28-3 (63-222°K; 20-8-6 X 107 Hz). kk . u Dielectric absorption gives com­ plex dielectric constants; be­ haviour very sensitive to phase transitions. 3-55-2-90 (222-6-295°K) n

HI

- 3 3 - 9 to - 3 2 - 9 (195-273°K) 00 +4·35 ρ ρ

- 2 3 - 6 to -22-6(195-273°K)°° -25« +0·45 ρ ρ + 13·25ρρ

-47·2(195°Κ) 0 0 - 4 8 - 3 to - 4 7 - 7 (233-281 °K)°°

Table of Atomic Weights, 1969, IUPAC Commission on Atomic Weights; Pure Appl. Chem. 21 (1970) 95. G. Herzberg, Molecular Spectra and Molecular Structure. I. Spectra of Diatomic Molecules, 2nd edn., pp. 534-540. Van Nostrand (1950). Gmelins Handbuch der Anorganischen Chemie, 8 Auflage, "Chlor", Teil B (1968). Tables of Constants and Numerical Data, No. 17, Spectroscopic Data relative to Diatomic Molecules (general ed. B. Rosen), Pergamon (1970). A. L. McClellan, Tables of Experimental Dipole Moments, Freeman, San Francisco (1963); F. A. Van Dijk and A. Dymanus, Chem. Phys. Letters, 5 (1970) 387. f S. Weiss and R. H. Cole, / . Chem. Phys. 46 (1967) 644. β D. H. Rank, U. Fink and T. A. Wiggins, / . Mol. Spectroscopy, 18 (1965) 170. h Mellofs Comprehensive Treatise on Inorganic and Theoretical Chemistry, Supplement II, Part I, Longmans, Green and Co. (1956). i _ S. C. Hurlock, R. M. Alexander, K. N. Rao, N. Dreska and L. A. Pugh, / . Mol. Spectroscopy, 37 (1971) 373. i National Bureau of Standards Technical Note 270-3, U.S. Govt. Printing Office, Washington (1968); Codata Bulletin, International Council of Scientific t£ <-* Unions, Committee on Data for Science and Technology (1970). k H. J. Lempka, T. R. Passmore and W. C. Price, Proc. Roy. Soc. A304 (1968) 53. 1 F. D. Rossini, D. D. Wagman, W. H. Evans, S. Levine and I. Jaffe, Selected Values of Chemical Thermodynamic Properties, Circular 500, National Bureau of Standards, Washington (1952). m G. Woolsey, / . Amer. Chem. Soc. 59 (1937) 1577. n J. R. Bates, J. O. Halford and L. C. Anderson, / . Chem. Phys. 3 (1935) 531. 0 E. Sändor and R. F. C. Farrow, Nature, 213 (1967) 171. p N. Niimura, K. Shimaoka, H. Motegi and S. Hoshino,/. Phys. Soc. Japan, 32 (1972) 1019. q E. Sändor and R. F. C. Farrow, Nature, 215 (1967) 1265. r E. Sändor and M. W. Johnson, Nature, 217 (1968) 541. 8 F. A. Mauer, C. J. Keffer, R. B. Reeves and D. W. Robinson, / . Chem. Phys. 42 (1965) 1465. * T. Tokuhiro, / . Chem. Phys. 47 (1967) 109. u (1) D. F. Hornig and W. E. Osberg, / . Chem. Phys. 23 (1955) 662; G. L. Hiebert and D. F. Hornig, ibid. 26 (1957) 1762; 27 (1957) 752, 1216; 28 (1958) 316; (2) A. Anderson, H. A. Gebbie and S. H. Walmsley, Mol. Phys. 7 (1964) 401; (3) L.-C. Brunei and M. Peyron, Compt. rend. 264C (1967) 821, 930; (4) R. Savoie and A. Anderson, / . Chem. Phys. 44 (1966) 548; (5) J. Blanchard, L.-C. Brunei and M. Peyron, Compt. rend. 272B (1971) 366; (6) M. Ito, M. Suzuki and T. Yokoyama, /. Chem. Phys. 50 (1969) 2949; (7) E. L. Pace, Spectrochim. Acta, 27A (1971) 491.

a b c d e

Diamagnetic susceptibility, X(x 106 c.g.s. u n its mol" 1 ) Solid Liquid Gas *H nmr chemical shift of gas (ppm relative to CH 4 )

Michigan (1960-8). A. Eucken and E. Schröder, Ann. Phys. 36(5) (1939) 609. * H. Ziebland and D. P. Needham, Proceedings of the 4th Symposium on Thermophysical Properties (ed. J. R. Moszynski), p. 296, American Society of Mechanical Engineers, New York (1968). ·» E. U. Franck, Z. Elektrochem. 55 (1951) 636. bb A. K. Barua, A. Manna and P. Mukhopadhyay, / . Chem. Phys. 49 (1968) 2422; C. E. Baker and R. S. Brokaw, ibid. 43 (1965) 3519. cc W. Biltz and A. Lemke, Z. anorg. Chem. 203 (1932) 321. dd S. J. Yosim, / . Chem. Phys. 40 (1964) 3069. ee J. W. Stewart, / . Chem. Phys. 36 (1962) 400. » W. G. Strunk and W. H. Wingate, / . Amer. Chem. Soc. 76 (1954) 1025. ee Gmelins Handbuch der Anorganischen Chemie, 8 Auflage, "Brom" (1931). hh Gmelins Handbuch der Anorganischen Chemie, 8 Auflage, "Iod" (1933). 11 Landolt-Börnstein Tables, II Band, Eigenschaften der Materie in ihren Aggregatzuständen, 5 Teil, Bandteil a, "Transportphänomene I" (1969). W Gmelins Handbuch der Anorganischen Chemie, 8 Auflage, "Chlor" (1927). kk P. P. M. Groenewegen and R. H. Cole, / . Chem. Phys. 46 (1967) 1069; R. H. Cole and S. Havriliak, jun., Discuss. Faraday Soc. 23 (1957) 31. 11 Landolt-Börnstein Tables, II Band, Eigenschaften der Materie in ihren Aggregatzuständen, 6 Teil, "Elektrische Eigenschaften I" (1959). 13X111 M. E. Peach and T. C. Waddington, Non-aqueous Solvent Systems (ed. T. C. Waddington), p. 85. Academic Press (1965). nn N. Bauer and K. Fajans, / . Amer. Chem. Soc. 64 (1942) 3023. 00 Landolt-Börnstein Tables, II Band, Eigenschaften der Materie in ihren Aggregatzuständen, 10 Teil, "Magnetische Eigenschaften II" (1967). pp J. A. Pople, W. G. Schneider and H. J. Bernstein, High-resolution Nuclear Magnetic Resonance, p. 90, McGraw-Hill (1959).

v W. F. Giauque and R. Wiebe, / . Amer. Chem. Soc. 50 (1928) 2193. w W. F. Giauque and R. Wiebe, /. Amer. Chem. Soc. 51 (1929) 1441. X JANAF Thermochemical Tables, The Dow Chemical Company, Midland, y

1297

THE HYDROGEN HALIDES

All the hydrogen halides are colourless. The gases fume strongly in moist air and have penetrating, acrid odours. In the absence of a polar solvent, the properties are exclusively those of molecular compounds, but dissolution in a solvent like water brings about dis­ sociation to give the solvated proton and halide ion. The hydrogen halides are thus acceptor molecules with a marked capacity, in Mulliken's terminology437, to form "inner com­ plexes" : S + HX

Polar solvent

^

S,HX solv

^

Outer complex

[SHl + X- 8 0 l v

Inner complex

This distinctive property is related to the relative proximity to the electronic ground state of excited states of the type H + X~ and to the large solvation energy accruing to the proton. By contrast with hydrogen fluoride, hydrogen chloride, bromide and iodide are all strong acids in aqueous solution. Table 29 also includes recommended values of spectroscopic properties, vapour pres­ sures, heat capacities and thermodynamic functions for the hydrogen halides. Satisfactory agreement is found between the thermodynamic functions derived via the third-law method from measurements of heat capacity and those calculated from spectroscopic data, so that there is no evidence that the solids retain residual entropy at limiting low temperatures. Whereas the standard entropies vary but little from one gaseous molecule to the other, the bond energy sequence HC1 > HBr > HI reflects the marked decline in — ΔϋΓ/. It is this enthalpy term which dictates the relative stabilities of the gaseous hydrogen halides with respect to thermal dissociation; the results of Table 30 demonstrate the dramatic TABLE 30. THERMAL STABILITY OF THE GASEOUS HYDROGEN HALIDES»

HC1 Temperature (°K) 100 200 300 500 1000 1500 2000 3000 4000 6000

% dissoc*

— —

8-4x10-4 0-23 1-53 3-89 9-58 14-7 220

HBr log Kp 48-709 24-624 16-596 10151 5-265 3-615 2-785 1-950 1-529 1097

% dissoc.*

— —

0-11 2-58 7-27 11*9 190 23-6 29-1

HI log Kp 19-520 12011 9-290 5-927 3-155 2-212 1-737 1-260 1-020 0-774

% dissoc.*

— —

22-9 301 33-1 34-6 361 36-9 37-8

logKp -9-920 -2-599 -0-247 + 1055 +0-732 +0-613 +0-554 +0-496 +0-466 +0-431

* At a pressure of 1 atm. *JANAF Thermochemical Tables, The Dow Chemical Company, Midland, Michigan (1960-8).

change of these stabilities, from hydrogen chloride, which is dissociated to the extent of no more than 9-6% even at 3000°K, to hydrogen iodide, 23% of which is dissociated at 500°K. Only the slowness of the decomposition, in the absence of a catalyst, keeps hydrogen iodide from being about 50% dissociated at room temperature. Details of the physical properties of the deuterium halides have not generally been in­ cluded in Table 29, but are to be found in references 425 and 426. 437 R . s . Mulliken and W. B. Person, Molecular Complexes* Wiley-Interscience, New York (1969).

1298

CHLORINE, BROMINE, IODINE AND ASTATINE: A. J. DOWNS AND C. J. ADAMS

2. Physical properties of the gaseous molecules. The 1Σ + electronic ground state of the hydrogen halide molecules involves the configuration (nsa)2(npa)2(npny for the valence electrons438. Molecular-orbital calculations indicate that the (nsa)2 and («/>π)4 electrons are mainly atomic and that the formation of the H-X bond is principally due to the (npo)2 electrons (see Fig. 24), though a small contribution probably derives from interaction with vacant hydrogen orbitals (e.g. 2ρπ). The electronic states ofthe molecules have been studied theoretically in great detail by Mulliken438. ηρσ

npss

-fj

(nsa)2

HX FIG. 24. Energy level diagram showing the molecular orbitals of a hydrogen halide molecule.

The electronic spectra ofthe molecules, studied in the ultraviolet region both in absorp­ tion and emission425»426»439, exhibit the following features: (i) a continuous absorption band at low frequencies attributed to the transition 3Π,1Π (ηρσ)2(ηρπ)3(ηρσ*) <-+ ι + Σ (ηρσ)2(ηρπ)4; (ii) banded systems represented as 3Π,1Π (npa)2(npny(n+l)sa<-+ Χ + Σ (ηρσ)2(ηρπ)49 where the excited states involve JJ-like coupling between the outer σ-electron and the H X + core; (iii) band systems ascribed to transitions of the type ιΣ + (npa)(npn)\n+\)sa <-+ιΣ + (ηρσ)2(ηρπ)4\ and (iv) bands arising from transitions involving low-lying states of the molecular ions HX + . An analysis of such measurements and ofthe photoelectron spectra439 ofthe molecules has led to the construction of potential energy diagrams such as that given in Fig. 25 for HC1. The disposition and form ofthe poten­ tial energy curves do not favour a banded spectrum at low frequencies analogous to that observed for Cl2, Br2 or I 2 . Accordingly, no accurate value for the dissociation energy of the HX molecule can be obtained directly from spectroscopic measurements. Instead, the most reliable value of this parameter, quoted in Table 29, is derived from thermochemical data. Reference to the electronic and photoelectron spectra associated with the molecular ions HX+ has provided the information about these species which is presented in Table 31. 438 R . s . Mulliken, Rev. Mod. Phys. 4 (1932) 1 ; Phys. Rev. 5 0 (1936) 1017; ibid. 5 1 (1937) 310. 439 H . J. L e m p k a , T . R . P a s s m o r e a n d W . C . Price, Proc. Roy. Soc. A304 (1968) 53 a n d references cited therein.

1299

THE HYDROGEN HALIDES

H+C1+

1

2

3 4 5 6 Internuclear distance, A

7

FIG. 25. Potential energy curves for the ground and excited states of HCl and HC1+

TABLE 31. PROPERTIES OF THE MOLECULAR IONS HX + »·*>

Molecular ion

HC1+

Electronic state

Internuclear distance» r,(A)

Vibrational frequency, AG*(cm"i)

kcal 108-3 106-5 41-5

eV 4-70 4-62 1-80

\

1-3153 1-5140

2568 1527

Πι/2

11-67 1200 15-29

90-6 82-9 47-5

3-93 3-60 206

\

1-448 1-684

2348 1329

Π 3 /2 2Π1/2 2Σ +

10-38 11 05 13-85

72-1 56-7 64-8

3-13 2-46 2-81

\

1-62 1-90

2170 1040

Π 3 /2

2

Πι/2

2

Π 3 /2

2

2Σ +

HI +

Dissociation energy,

12-74 12-82 16-23

2

2Σ +

HBr+

Adiabatic ionization potential relative t o ^ * ground state of HX(eV)

2

a Tables of Constants and Numerical Data, No. 17, Spectroscopic Data relative to Diatomic Molecules (general ed. B. Rosen), Pergamon (1970). b H. J. Lempka, T. R. Passmore and W. C. Price, Proc. Roy. Soc. A304 (1968) 53.

1300

CHLORINE, BROMINE, IODINE AND ASTATINE! A. J. DOWNS AND C. J. ADAMS

It is noteworthy that the internuclear distance is somewhat longer, and the vibrational fre­ quency somewhat lower, for the molecular ion in its ground state than for the parent molecule. Despite this evidence of weaker binding in the molecular ion, the energy of dis­ sociation of the ion is slightly greater than that of the molecule. However, the general tenor of the results gives clear notice that thepn electrons of the HX molecule, which are the first to ionize, are virtually non-bonding. The ionization energies of the ρπ electrons bear a linear relationship to the ionization energies of the corresponding noble gases, whence the hydrogen halides are formally derived by withdrawing a proton from the nucleus439. Detailed analyses of the vibrational-rotational and microwave spectra of the gaseous hydrogen halides yield the most reliable estimates of the internuclear distances re (see Table 29). Recent measurements of the Stark effect in the microwave spectra of hydrogen bromide and iodide give precise values for the dipole moments of these molecules440; with the corresponding parameters recommended on the basis of earlier measurements for the other hydrogen halides, these confirm the striking decrease in dipole moment along the series HF, HC1, HBr, HI. Quadrupole coupling constants have also been obtained from the microwave spectra of hydrogen, deuterium and tritium halides441. The significance of these parameters has been discussed in subsection 3.2 (pp. 1271-4), and estimates of the ionic character of the H-Br and H-I bonds, based on the classical Townes-Dailey approxi­ mation387, are included in Table 27. Nmr measurements on the gaseous hydrogen halides show that the proton resonance moves progressively to higher field in the sequence HCl < HBr < HI 374 ; the increased shielding of the proton thus implied is consistent with the variations of electron density to be expected as the electronegativity of the halogen decreases. The infrared and Raman spectra of the gaseous hydrogen halides give the values listed in Table 29 for the vibrational frequencies ωβ, the anharmonicity constants ωβχβ, and the stretching force constants ke. The influence of change of physical state on the infrared and Raman spectra of the hydrogen halides has been the subject of several investigations, though a full interpretation of the results has not always been possible356»426. The lowfrequency shift of the vibrational fundamental on going from the gaseous to the condensed phases provides clear evidence of intermodular interaction involving the hydrogen halide molecule, notably in the solid phase (q.v.) or when dissolved in solvents which do not promote dissociation. The infrared spectra of hydrogen halide molecules isolated in a variety of inert matrices at 20°K signify that diffusion occasions the formation of cyclic dimers and trimers and various multimeric species; in the presence of N 2 , C0 2 , CO or C2H4 (D) in argon matrices at 20°K, evidence is found for interactions of the type HX-D, HX-D-HX and (HX) 2 -D, while the HX-C2H4 system is believed to give rise to a specific complex442·443. 3. The solid phase. The solid hydrogen halides are noteworthy for their polymor­ phism. Thus, solid hydrogen chloride undergoes a first-order transition at 98-4°K; solid hydrogen bromide undergoes three λ-type transitions at 89-75°, 113-62° and 116-86°K, while for solid hydrogen iodide transition points have been detected at ca. 25°, 70° and 440 F . 441 T. 442 A . 443 A .

A . Van Dijk and A . Dymanus, Chem. Phys. Letters, 5 (1970) 387. Tokuhiro, / . Chem. Phys. 47 (1967) 109. J. Barnes and H . E. Hallam, Quart. Rev. Chem. Soc. 23 (1969) 398. J. Barnes, H . E. Hallam and G. F. Scrimshaw, Trans. Faraday Soc. 65 (1969) 3150, 3159, 3172.

THE HYDROGEN HALIDES

1301

444

125°K . Of the different forms, those stable at low temperatures have been most fully characterized. On the evidence of neutron diffraction studies of deuterium chloride445 and bromide446 at low temperatures, the crystals adopt a face-centred orthorhombic lattice with a unit cell belonging to the space group Bb2\m and containing four molecules, the dimen­ sions being as listed in Table 29. Within the generous limits of error, the intramolecular H-X distance is not significantly different from that of the gaseous molecule. The structure of linear zigzag chains of HX molecules is confirmed by the vibrational spectra of the solids 447 " 450 , though recent interpretations449»450 of these spectra favour non-planar rather than the planar chains implicit in the space group of the unit cell deduced by neutron diffraction. Qualitatively, therefore, the structure resembles that found in gaseous and crystalline hydrogen fluoride, and there are several reasons for believing that it is based, not on a simple dipolar array, but on hydrogen-bonding between adjacent HX molecules. Such hydrogen-bonding is, however, relatively weak. At low temperatures hydrogen iodide differs from the chloride and bromide in adopting an ordered face-centred tetragonal lattice451. At higher temperatures, all the solids appear to assume disordered face-centred cubic lattices, again with four molecules per unit cell426»451»452. A model has been proposed452 for this phase of deuterium chloride, in which the deuterium atom has twelve equally prob­ able equilibrium positions situated along the lines connecting the chlorine atom of the DC1 molecule with its twelve nearest neighbours; the reorientational motions of the hydrogen halide molecules have also been discussed450 in relation to the observed Raman spectra of the high-temperature phases. The relative mobility of molecules in the solid hydrogen halides is also indicated by the ferroelectric character reported for the orthorhombic phase of hydrogen chloride453, and by the complicated picture presented by measurements of the dielectric properties of the solids454. Properties of the Liquids: Function as Non-aqueous Solvents432»436»455 The anhydrous hydrogen halides HC1, HBr and HI have some interest as solvent systems, both in their own right and also because of the comparison thus allowed with liquid hydro­ gen fluoride (see Chapter 25)456. The first experiments using liquid hydrogen chloride as a solvent were reported in 1865 by Gore457, who, on the basis of qualitative visual observations of solubility, concluded that the liquid "has but a feeble solvent power for solid bodies in 444 G . L. Hiebert and D . F . Hornig, / . Chem. Phys. 26 (1957) 1762. 445 E . Sändor and R. F . C. Farrow, Nature, 213 (1967) 171. 446 E . Sändor a n d M . W . Johnson, Nature, 217 (1968) 541. 447 D . F . Hornig a n d W . E . Osberg, / . Chem. Phys. 23 (1955) 662; G . L . Hiebert a n d D . F . H o r n i g , ibid. 27 (1957) 752, 1216; ibid. 28 (1958) 316. 448 A . Anderson, H . A . Gebbie a n d S. H . Walmsley, Mol. Phys. 7 (1964) 4 0 1 . 449 R . Savoie a n d A . Anderson, / . Chem. Phys. 4 4 (1966) 5 4 8 ; L . - C . Brunei a n d M . Peyron, Compt. rend. 2 6 4 C ( 1 9 6 7 ) 8 2 1 . 450 M . I t o , M . Suzuki a n d T . Yokoyama, / . Chem. Phys. 50 (1969) 2949. 451 F . A . M a u e r , C . J . Keffer, R . B . Reeves a n d D . W . Robinson, / . Chem. Phys. 42 (1965) 1465. 452 E . Sändor a n d R . F . C . F a r r o w , Nature, 215 (1967) 1265. 453 K . Shimaoka, N . Niimura a n d S. Hoshino, Acta Cryst. A25 (1969) S50. 454 p . p . M . Groenewegen a n d R . H . Cole, / . Chem. Phys. 46 (1967) 1069; R . H . Cole a n d S. Havriliak, jun., Discuss. Faraday Soc. 23 (1957) 31. 455 T . c . Waddington, Non-aqueous Solvents, Nelson, L o n d o n (1969); R . A . Zingaro, Nonaqueous Solvents, Raytheon Educ. C o . , Boston (1968). 456 H . H . H y m a n a n d J. J. K a t z , Non-aqueous Solvent Systems (ed. T . C . Waddington), p . 47. Academic Press (1965). 457 G . G o r e , Phil. Mag. 29 (4) (1865) 541.

1302

CHLORINE, BROMINE, IODINE AND ASTATINE: A. J. DOWNS AND C. J. ADAMS

general". Nevertheless, the study of hydrogen chloride, bromide and iodide as non-aqueous solvents was resumed just after the turn of the century when Steele, Mclntosh and Archibald described physical and thermodynamic constants of the liquids, as well as solubility, conduc­ tivity, transport number and ebullioscopic measurements458; little correlation could be found between dissolving power, dielectric constant, conductivity and the molecular weights measured for solute species. More recently, there have been extensive studies of the liquids, particularly hydrogen chloride, as reaction media432. Of the physical properties of the hydrogen halides recorded in Table 29, the relatively narrow liquid range (15-30°), with its implied experimental difficulties, and low dielectric constant (3-14) are of especial significance in relation to the exploitation of the liquids as solvents. By contrast, hydrogen fluoride has a liquid range of 100° and a dielectric constant of 175 at — 73°C. The low dielectric constants of hydrogen chloride, bromide and iodide mean that only salts with low lattice energies, e.g. the tetra-alkylammonium halides, are appreciably soluble, and that extensive ion-pairing occurs in all but the most dilute solutions. The electrical conductivities of the liquids, though lower than those of hydrogen fluoride and water, suggest some form of self-ionization such as 3HX^—H2X++HX2-

with solvation of both the proton and halide ion produced by the primary dissociation pro­ cess. The only reasonably well authenticated example of an H2X + ion in solution is that where X = F, which has been detected in the HF-SbF5 system. The postulated H2CI+ ion is known to exist in the gas phase459; it may also be present in the compounds HC1,HC104, HC1,H2S04 or HCl,HBr, though there is no convincing evidence to this effect. The ions H2Br + and H 2 I + have likewise eluded detection in the condensed phases. On the other hand, the existence of the hydrogen dihalide anions HX2~ is now well estab­ lished (see pp. 1313-21); the presence of more extensively solvated halide ions, e.g. Cl(HCl)n~ (n > 1), has also been considered460. If the self-ionization is correctly represented by the above equation, two definitions of acids and bases are applicable, based upon a difference of emphasis rather than of principle; this arises from the fact that either halide ion- or proton-transfer can be regarded as the primary step in the equilibrium. Accordingly, acids can be defined as either proton-donors or halide ion-acceptors, and bases as either protonacceptors or halide ion-donors. Thus, hydrogen chloride shows some of the characteristics of a "chloridotropic" solvent like arsenic trichloride and some of an acidic solvent like sulphuric acid. Because of the highly acidic nature of the solvents, it has been impossible to find any reasonably strong solvo-acids, though the following compounds appear to function as weak acids: boron(III) halides, tin(IV) halides, and, in liquid hydrogen chloride, PF5 (the strongest acid yet found), XCl and HX (X = Br or I). Conversely, there are very many solvo-bases, some of which act as strong electrolytes. Such bases comprise (a) salts which ionize readily to give a free halide ion, e.g. tetra-alkylammonium halides, and (b) compounds containing atoms with a lone pair of electrons, e.g. amines, phosphines, ethers 458 B. D . Steele, D . Mclntosh and E . H . Archibald, Phil. Trans. 205 (1905) 9 9 ; Z. physik. Chem. 55 (1906) 129. 459 D . O. Schissler and D . P. Stevenson, / . Chem. Phys. 2 4 (1956) 926; F . H . Field and F . W. Lampe, / . Amer. Chem. Soc. 80 (1958) 5583. 4öo C. J. Ludman and T. C. Waddington, 2nd International Raman Conference, Oxford (1970); T. C. Waddington, 6th International Symposium on Fluorine Chemistry, Durham (1971).

THE HYDROGEN HALIDES

1303

or sulphides, or a π-bonded system which can easily be protonated, e.g. aromatic olefins and organic compounds containing the groups - C = N , - N = N - , N c = 0 or ->P=0. Attention has been directed, not only to acid-base behaviour, but also to solvolysis and redox reactions in liquid hydrogen halide solutions432. Solvolysis corresponds to the replacement of a ligand Y by a halogen atom X, a type of reaction which has been ob­ served432·461 in solutions in hydrogen chloride when Y is phenyl, hydroxyl or fluorine, e.g. Ph3SnCl+HCl -> Ph2SnCl2+PhH Ph3COH+3HCl -> Ph3C+HCl2- + H 3 0 + C1SbF3+3HC1 -> SbCl3+3HF

By means of conductimetric titrations, it has been possible to establish, inter alia, the follow­ ing oxidation reactions of chlorine, bromine and iodine monochloride in liquid hydrogen chloride432: i-+ 2Ci2— i c i Br -+Cl2-— BrClJ I" +

.„. IC1 +

...-,. HC1

oxidation w , -IU + IC1 acid-base f

HC1 + IC12PC13 4-

X2 +

HC1

*· PCI3X++ HC1X- (X = Cl or Br)

Experimental studies of the solvent properties of the hydrogen halides normally demand that the liquids be handled at low temperatures in an enclosed vacuum system. The methods principally exploited to investigate the behaviour of solutes and to monitor reactions in solution are as follows. Conductimetric Measurements These have been the mainstay of many recent investigations432. The mechanism of the conduction of acids and bases is not known, but probably involves halide ion-transfer, at least in basic solution, e.g. ci—H— cr

H—ci

\ Cl

H

Cl—H—Cl-

The variation of the molar conductance Am with the concentration c of strongly basic solutions is noteworthy, for plots of log Am against log c or of Am against Vc tend to exhibit a minimum, a behaviour characteristic of solvents of low dielectric constant432. To account for these results, two modes of ionization have been suggested for an ion-pair A + B - : A+B-

A+B- + A+ AHB- +

B-

^=^ A+ + B- very dilute solutions more concentrated ^=^ A2B+ Λ solutions *. ABi J

The variation of the conductance of weak electrolytes is more complicated, and cannot 461 M. E. Peach, Inorg. Nuclear Chem. Letters, 7 (1971) 75.

1304

CHLORINE, BROMINE, IODINE AND ASTATINE: A. J. DOWNS AND C. J. ADAMS

easily be explained. Reactions involving changes in the ionic species present in solution can be monitored conductimetrically. The processes of salt-formation, adduct-formation and ionization, or a combination of these, give characteristic conductimetric curves. Hence, it has been possible to follow the course of acid-base reactions such as Me 4 N + HCl 2 - +BCI3 -> Me 4 N + BCl 4 - +HC1

or of redox reactions such as PCl3 + Cl2+HCl->PCl4 + HCl 2 -

Spectroscopic Measurements The spectroscopic properties of solutions in hydrogen halides have been very little studied, mainly because of the technical difficulties of maintaining the sample either at low temperatures or at relatively high pressures. However, nmr spectra have confirmed the presence of Ph 2 CCH3 + in solutions of Ph 2 C = CH 2 and of PhCCH3 2 + in solutions of PhC=CH, both in liquid hydrogen chloride at room temperature432. The weakness and simplicity of the Raman spectrum of the pure liquid facilitate the application of Raman spectroscopy to solutions in hydrogen chloride; according to preliminary reports460, such spectra serve to identify certain solute species in solutions at room temperature. Phase Diagrams While not revealing much about the nature of the solutions, phase diagrams in which one of the hydrogen halides is a component give indications of some of the compounds that may be formed in solution, and may also yield valuable information about the function of solvo-acids and solvo-bases. Thus, the fact that the system BCI3-HCI gives no sign of compound-formation suggests that BC13 must be a very weak solvo-acid in hydrogen chloride. The diagrams for several compounds that are solvo-bases with hydrogen chloride or bromide have also been studied. Cryoscopic and Ebullioscopic Measurements These have been made on some solutions in hydrogen chloride, bromide and iodide, but the results are commonly difficult to interpret because of the complex mode of ionization432. Preparative Methods Because of their low boiling points and consequent easy removal, the liquid hydrogen halides are useful media for certain preparations. Convenient routes to the following species or their derivatives have thus been devised: BX4~ 432, BF 3 C1- 432 , B 2 C1 6 2 - 432 , N02C1432, A12C17 - 462 , R2SC1 + and RSC12 + 4 « , PCl3Br + 432, Ni 2 Cl 4 (CO) 3 and Ni(NO)2Cl2 (products of the reactions of Ni(CO)4 with Cl 2 and NOC1, respectively)464, [Fe(CO)5H]+ and [(7r-C5H5)Fe(CO)2]2H+ and [(^-C5H5)Fe(CO)]4H22+ (formed by protonation of the appro­ priate neutral species) as well as the cation [(7r-C5H5)Fe(CO)]4+465. Solutions of the Hydrogen Halides In common with hydrogen fluoride, the heavier hydrogen halides are notable for their profuse solubility in water; in no case do the solutions comply even approximately with Henry's law. Details of the solubility and of the density of the resulting solutions are pre­ sented in Table 32. Each of the aqueous systems gives rise to a maximum-boiling azeotrope, 462 M . E . Peach, V. L . Tracy and T. C . Waddington, / . Chem. Soc. {A) (1969) 366. 463 M . E . Peach, Canad. J. Chem, 4 7 (1969) 1675. 464 z . Iqbal and T. C . Waddington, / . Chem. Soc. (A) (1968) 2958; ibid. (1969) 1092. 465 D . A. Symon and T. C. Waddington, / . Chem. Soc. {A) (1971) 953; / . Chem. Soc.y Dalton Trans. (1973) 1879.

HX(aq) -> H+(aq) + X~(aq), Δβ° calc.

112 38 1-138

1035 1072

Miscible in all proportions

HF

+

4-85

+ 5-65 + 128-6 +232-9 -362-3

undissociated ionized -11-70 -14-70 - 5-65 - 1-34 -20-3 -44-8

Free energy changes for ionization of HX molecules in water at 25°C, AG° (kcal mol"*) 1. HX(aq)->HX(g) 2. HX(g)-+H(g) + +X(g) 3. H(g) + X(g)->H (g) + X-(g) 4. H+(g) + X-(g)-*H + (aq) + X~(aq)

Thermodynamic functions for the process HX(g) + ooH20->H+X-(ooH20) at 25°C: AJTGccalmori) AG°(kcalmol-i) AS°(caldeg-imol-i)

Constant-boiling solution at 1 atm boiling point (°C) concentration (g/100 g solution) density (gcm-3 at 25°C)

Density of aqueous solution (gem - 3 at 20°C) concentration: 10 g/100 g solution 20 g/100 g solution saturated

Solubility in water (g/100 g soln at 1 atm)

Property

-

9-5

- 1-2 + 96-6 +229-1 -3340

-17-890 - 8-595 -31-1 5

108-58 20-22 1096

1047 1091 1-205

4515 at 0°C 4202 at 20°C 37-34 at 50°C

HC1

- 11-9

- 0-7 + 81 -05 +234-9 -327-1

-20-35 -1208 -27-8

124-3 47-63 1-482

1073 1-158 1-79

68-85 at 0°C 65-88 at 25°C 63-16 at 50°C

HBr

TABLE 32. BEHAVIOUR OF THE HYDROGEN HALIDES IN AQUEOUS SOLUTION* ~ f

at 0°C

- 11-9

- 0-5 + 650 +241-8 -318-2

-19-52 -12-74 -22-7 5

126-7 56-7 1-708

1072 1167 1-99

~71

HI

-

-

0-5 18 3-6 3-2

HF

~_

14 70 7

13-7

HC1

<-

11 8-8 7

15-2

HBr

<-

-

7-4 8-8 7

141

HI

a J. W. Mellor, A Comprehensive Treatise on Inorganic and Theoretical Chemistry, Vol. II, Longmans, Green and Co., London (1922); Supplement II, Part I (1956). b Gmelins Handbuch der Anorganischen Chemie, 8 Auflage, "Chlor" (1927); "Brom" (1931); "Iod" (1933). c Z. E. Jolles (ed.), Bromine and its Compounds, Benn, London (1966). d A. G. Sharpe, Halogen Chemistry (ed. V. Gutmann), Vol. 1, p. 1. Academic Press (1967). e J. C. McCoubrey, Trans. Faraday Soc. 51 (1955) 743. f National Bureau of Standards Technical Note 270-3, U.S. Govt. Printing Office, Washington (1968).

HX(aq) -> H + (aq) + X"(aq), AH° calc. (kcal mol - 1 ) HX(aq)->H + (aq) + X"(aq), AS° calc. (cal deg"1 mol" 1 ) Calculated pK a for acid = AG°/l-36 Observed pKa for acid

Property

TABLE 32 (cont.)

THE HYDROGEN HALIDES

1307

the characteristics of which are listed in the table. The stability of composition provided by constant-boiling hydrochloric acid has led to its widespread use as a standard for ana­ lytical purposes; samples of the constant-boiling acid prepared by distillation have been found, after storage for more than three years, to differ in composition by less than 0-1% from freshly prepared samples426. Unlike hydrogen fluoride, the other hydrogen halides undergo virtually complete ionization in aqueous solution at all but the highest concentrations. Thermodynamic functions defining the dissolution and dissociation of the hydrogen halides in water are included in Table 32, wherein the thermodynamic parameters for the process HX(aq) ^ H

+

(aq)+X"(aq)

are related to the sum of the appropriate parameters for the following stages: 1.

HX(aq)->HX(g)297.466

2. 3.

HX(g)->H(g)+X(g)297 H(g)+X(g) - * H + (g)+X"(g) 2 8 9 · 2 9 7 +

4. H (g)+X-(g)->H + (aq)+X-(aq)296

For reactions (1), (2) and (4), free energy and enthalpy changes have been derived from published measurements or estimates; for reaction (3), the free energy has been estimated from the heats of ionization of the hydrogen and halogen atoms289»297, together with an entropy contribution of -R loge 4—Ä loge 2 = -4-13 cal deg - 1 mol -1 , due to the change of electronic multiplicity. Such an analysis reproduces, within the limits of error, the observed pattern of ionization, even though a relatively small change of free energy, namely < 10 kcal, is implied by the distinction between a strong acid like HC1 and a weak acid like HF. The results demonstrate that, in acid strength, there is little to choose between aqueous hydrochloric, hydrobromic and hydriodic acids. The difference between hydrogen fluoride and the other hydrogen halides resides mainly in the more endothermic enthalpy of ionization; this, in turn, arises primarily from the higher bond energy of the HF molecule, but partly also from the enhanced hydration energy of the undissociated HF molecule (the consequence of hydrogen-bonding to water molecules) and the reduced electron affinity of fluorine. These factors together more than compensate for the high hydration energy of the fluoride ion. The physical properties of the aqueous hydrohalic acids, and particularly hydrochloric acid, have been the subject of many investigations. Included in more comprehensive accounts of the halogens418»424 ~426 are details of the following properties for one or more of the aqueous hydrohalic acids: density, vapour pressure, specific heat, heats of solution and neutralization, activity coefficients, viscosity, surface tension, compressibility, diffusion coefficients, distribution coefficients between water and other liquids, molal volume, electrical conductivity, transport numbers of the ions, dielectric constant, refrac­ tive index and magnetic susceptibility. Cooling aqueous solutions of the hydrogen halides produces a variety of solid hydrates, the properties of which are summarized in Table 33. The vibrational spectra and recent X-ray analyses leave little doubt that these hydrates are to be formulated as [(H 2 0) n H] + X _ . The compounds are of interest in the opportunity they provide for the study of the hydrated 4 6

6 J. C. McCoubrey, Trans. Faraday Soc. 51 (1955) 743.

X = Brb.c X= P

Very unstable, m.p. -70°C.

HX,6H20

H

H

.O-H-a

HX,4H20

«*■·■

M.p. -24-9°C. Shown by X-ray analysis to be H 5 0 2 + C1-,H 2 0 with nearly eclipsed (H 2 0) 2 H +

.O-H—O' H"/ ^H H Central O-H distances = 2-41Ä.

-H +

M.p. -17-7°C. Shown by X-ray analysis to be (H20)2H+C1~; the bonding arrangement around one end of the (H 20)2H+ ion is almost planar and pyra­ midal around the other.'

M.p. ca. -48°C.

M.p. ca. -43°C.

Decomposes -88-2°C.

M.p. — 55-8°C. Shown by X-ray analysis to be M.p. -36-5°C. [(H20)3H]+[(H20)4H]+2Br-,H20 with H-O .bond lengths of 2-465-2-75 A. The aggregates [(H20)3H]+ and [(H20)4H]+ may be regarded as H 3 0 + ,2H 2 0 and H 3 0 + ,3H 2 0, respectively.11

Decomposes -47-9°C.

M.p. -11-3°C.

spectrum M.p. — 15-35°C. Indicated by i.r. spectrum0 and con­ Stable between —3-3° and — 15-5°C under pressure. I.r. consistent with firmed by X-ray analysis to be H30+C1" with I.r. spectrum consistent with the formulation the formula­ O-H · Cl = 2-95A and O-H = 0-96 ± 008Ä(see H 3 0 + Br-. d tion H 3 0 + I~. d Fig. 26).e

X = Cla

Central O-H distahces = 2-43 A. Not known.

ΗΧ,3Η20

HX,2H20

HX,H 2 O

Formula

TABLE 33. SOLID HYDRATES OF THE HYDROGEN HALIDES

Gmelins Handbuch der Anorganischen Chemie, 8 Auflage, System-nummer 6, "Chlor", Teil B, Lieferung 1 (1968). Z. E. JoUes (ed.), Bromine and its Compounds, Benn, London (1966). J. W. Mellor, A Comprehensive Treatise on Inorganic and Theoretical Chemistry, Vol. II, Longmans, Green and Co., London (1922). C. C. Ferriso and D. F. Hornig, /. Chem. Phys. 23 (1955) 1464. Y. K. Yoon and G. B. Carpenter, Acta Cryst. 12 (1959) 17. J.-O. Lundgren and I. Olovsson, Acta Cryst. 23 (1967) 966. * J.-O. Lundgren and I. Olovsson, Acta Cryst. 23 (1967) 971. h J.-O. Lundgren and I. Olovsson, / . Chem. Phys. 49 (1968) 1068.

a b c d e 1

1310

CHLORINE, BROMINE, IODINE AND ASTATINE: A. J. DOWNS AND C. J. ADAMS

proton. Thus, X-ray analysis of single crystals at low temperatures has identified the aggre­ gates H3O+, [(H20)2H]+, [(H 2 0) 3 H]+ and [(H 2 0) 4 H] + 467, while some of the earliest direct evidence of the Η3<3+ ion was gained from the infrared spectra of the monohydrates at -195°C468.

FIG. 26. Crystal structure of HC1,H 2 0: (a) [111] plane; (b) [121] plane. [Reproduced with permission from Y. K. Yoon and G. B. Carpenter, Acta Cryst. 12 (1959) 20.]

Solutions of the hydrogen halides in non-aqueous solvents have also been the focus of some attention; notwithstanding the dearth of systematic measurements of solubility, a significant weight of information, both qualitative and quantitative, has accumu­ lated418»424 -426,469. Though limited by the reactivity of the solute, a wide range of solvents has thus been identified, extending from non-polar to highly polar media, e.g. hydrocarbons, carbon tetrachloride, chloroform and other organo-halogen compounds, alcohols, diethyl ether, sulphur dioxide, pyridine, aniline, N-methylacetamide, acetonitrile, nitrobenzene, acetic acid and hydrogen fluoride. The form of the solute varies correspondingly from HX molecules to solvated H + and X~ ions, as may be judged by the electrical conductivity. It has been concluded that the tendency of the hydrogen halides to dissociate in solution is largely determined by the capacity of the solvent molecule to act as a proton acceptor; the dielectric constant of the medium is thought to have a much smaller effect. Heats of solution of hydrogen chloride and hydrogen bromide in solvents like carbon tetrachloride, chloro­ form and various hydrocarbons297, falling typically in the range — 2-6 to —4-2 kcal, are relatively more exothermic than those of the parent halogens. The vibrational spectra of solutions of the hydrogen halides, like those of the parent halogens, in organic solvents show variations attributable to specific interactions between the solute and solvent mole­ cules;; the extent of hydrogen-bonding between HX and solvent molecules has thus been gauged in a qualitative manner. The very sparing electrical conductivity of many solutions, e.g. in benzene or sulphur dioxide, is greatly increased by the presence of traces of water, 467 Y . K . Yoon and G. B. Carpenter, Acta Cryst. 12 (1959) 17; J.-O. Lundgren and I. Olovsson, / . Chem. Phys. 49 (1968) 1068. 468 c . C. Ferriso and D. F. Hornig, / . Chem. Phys. 23 (1955) 1464. 469 w . F. Linke, Solubilities: Inorganic and Metal-organic Compounds, 4th edn., Vol. 1, van Nostrand, Princeton (1958).

THE HYDROGEN HALIDES

as a result of the reaction

1311

HX+H 2 0-*H 3 0 + +X-

In an acidic solvent like anhydrous acetic acid, both hydrogen chloride and hydrogen bromide behave as weak acids, the pK values of 8-85 and 6-40, respectively, giving clear notice that hydrogen bromide is the stronger acid under these conditions426. The order of acid strengths HI > HBr > HC1 is reported to prevail in acetonitrile, various alcohols, pyridine and aniline. Chemical Behaviour^-426,436

The hydrogen halide molecules share with diatomic halogen or interhalogen molecules the primary function of electron acceptors. However, relative to the ground state, inter­ action of the HX molecule with a donor species D, like water, ammonia or an organic base, has a markedly greater stabilizing influence on excited states represented approximately by the formulation H + X -, as a result of which the "outer complex" initially formed is very much more prone to transformation into an "inner complex": HX+D ^ D, HX jr-* [DH]+X" outer complex

inner complex

the net result being the heterolytic fission of the H-X bond. This mechanism is character­ istic of many solution reactions of the hydrogen halides, which, in their action as protondonors, behave as acids in the more restricted sense defined by Lowry and Bronsted; it is as proton-donors, for example, that the hydrogen halides react with metals and metal oxides, hydroxides, carbonates and sulphides. The halide ions released by heterolytic fission may suffer various possible fates: thus, they may be stabilized as such by solvation or by incor­ poration in a solid lattice; they may give rise to complex ions, e.g. I 3 ~, GaBr4 ~ or PtCl62 ~; or they may undergo oxidation to the parent halogen or even, in some circumstances, to oxy-halogen species like IO3 ~. Homolytic fission of the HX molecule HX->H+X

which occurs with increasing readiness in the series HC1 < HBr < HI, represents an alternative mechanism for reaction, being favoured at elevated temperatures in the gas phase or in solution in non-polar solvents. This process is facilitated by photolysis or radiolysis or by the agency of suitable catalysts. Because of these different mechanisms, with their dependence on the reaction medium, the chemical behaviour of the hydrogen halides is unusually sensitive to the nature of this medium. Thus, the kinetic barrier to both homolytic and heterolytic fission is such for the anhydrous materials that they are relatively inert, e.g. with respect to most metals and their oxides. By contrast, the reactivity of solutions in highly polar media like water has for long been a familiar feature of inorganic chemistry. The reactions of the hydrogen halides may also be classified, according to their outcome, as either addition or substitution. Addition reactions may be further sub-divided into those wherein the HX bond remains intact, as in the formation of hydrogen dihalide anions, and those wherein this bond is broken, as when a hydrogen halide adds to the ; C = C ^ or

1312

CHLORINE, BROMINE, IODINE AND ASTATINE: A. J. DOWNS AND C. J. ADAMS

— G = E C — units of organic molecules. Substitution reactions are exemplified by protonation, e.g. M + /1H3O

and

+

+ « X - -> M" + + «X + /1H2O + w/2H2

02-+2H30++2X-

by oxidation, e.g. and by exchange, e.g.

^2H20+2X-

X0 3 -+6H 3 0 + +5X- ->3X2+9H20 HX+D 2 ^DX+HD

Reactions representative of addition and substitution processes are also indicated sche­ matically in the flowchart of Fig. 27. HX I I -c-cI I

EX n + HY

-M-X + RH Exchange EY,

RX + R OH

ΧΗΥ - * Stabilized by large univalent cation

u Λ/ H X

Anion Y

H,+ X organic base B [M"H4]+X-or [BH]+X-

Mctal

chalcogenide Ϊ ^

ProtonA W'T» \ M H4 ation B5H8-, B 8 H 12 etc.

^w v , *· Μ Χ 2 η + HΗ O^

MXn+H2

M'HnX4.n + H2 B 5 H 9 , B 8 H 14 etc.

Key B =organic base, e.g. C 5 H 5 N or Me 3 N R = organic group R'= H or organic group E =various inorganic or organic species X or Y=: halogen M=metal atom M=Si or Ge M = N. P or As Q=0, S,SeorTe

χ - typically C l O ^ BrOj; IO7, ΜηΟ^, S,0 8 2 rvanadate or S e O j -

FIG. 27. Some representative reactions of the hydrogen halides.

1. Addition reactions in which the H-X bond remains intact. Numerous derivatives of the hydrogen halides are known in which it is likely, though seldom certain, that the H-X bond survives addition. Such compounds are formally analogous to the chargetransfer complexes of the molecular halogens, but, with the exception of anionic species

THE HYDROGEN HALIDES

1313

of the type HXY~ (see below), they remain relatively ill-characterized. Detailed studies of numerous HX-solvent systems reveal the formation of distinct compounds, which are commonly low-melting and stable with respect to dissociation only at low temperatures, though salt-like derivatives containing anions of the type X(HX)n ~ (see below) of appreciable stability may result from the interaction of hydrogen halides with certain organic bases. Spectroscopic measurements on the systems Me20,HCl47° and R3P,HX (R = Me or Ph; X = Cl, Br or I)471 signify the presence of a hydrogen-bonded molecular species, as distinct from ions. By contrast, materials most realistically represented by an ionic com­ position, e.g. [MeCONH3]+Cl-, [Et2OH] +[Cl(HCl)n] - or [MeC=NH] +[Cl(HCl)n] - 432, are formed by various ethers, nitriles and other organic bases susceptible to protonation; heterolytic cleavage of at least some H-X bonds is clearly implicit in such formulations. Neutral addition compounds which probably belong to the molecular category include HCl,HBr (somewhat implausibly formulated as [H2Cl]+Br-)426> 2H2S,3HBr345, Me2O,HC1470, olefin,wHCl (n = 1 or 2)472, acetylene,«HCl (« = 1, 2 or 4)472, Ar,HCl (Ar = aromatic hydrocarbon)472, R3P,HX47i, and clathrate compounds with phenols*4*, though the precise nature of many of these systems remains obscure. Addition of a halogen atom to a hydrogen halide molecule gives rise to a free radical short-lived under normal conditions. However, evidence has been obtained for the forma­ tion of the complex I · · · I-H on photolysis of ethyl iodide trapped in a hydrocarbon matrix at 77°K473. By contrast, the action of a discharge on a gaseous mixture of a hydrogen halide HX with the corresponding molecular halogen X 2 is believed to induce the reaction X+H-X -* X-H-X (X = Cl, Br or I) According to the infrared spectrum attributed by Pimentel, Noble and Bondybey to the HX 2 radical in the matrix-isolated condition, the stretching force constant of the H-X bond sees little change with the transition HX2~ -> HX 2 +e. Although this finding gives appealing support to the non-bonding character of the highest occupied molecular orbital of the HX2 " ion (see below), the conclusions are clouded by circumstantial evidence pointing to the possibility that the trapped species is actually HX2 - and not HX2474. Hydrogen Dihalide Anions and Related Species4*6*415 All the hydrogen halides HX have the capacity to function as acceptors with respect to a univalent anionic donor Y - , which may be a halide, pseudohalide or oxyanion, and so form relatively well-defined species of the type YHX ~. It is in the affinity of these species to trihalide ions like IC12~ that the analogy between the hydrogen halide and diatomic halogen or interhalogen molecules as acceptor species is most clearly manifest. The exist­ ence of the hydrogen dihalide anions requires hydrogen-bridging to exercise a primary rather than the secondary bonding function evident, for example, in the stabilization of crystal structures and in the association of liquids like ammonia, water and hydrogen fluoride476,477. However, the anions formed by the heavier halogens are notably less stable 470 G. Govil, A. D . H. Clague and H. J. Bernstein, / . Chem. Phys. 49 (1968) 2821. 47i M. Van den Akker and F. Jellinek, Rec. Trav. Chim. 86 (1967) 275. 472 D . Cook, Y. Lupien and W. G. Schneider, Canad. J. Chem. 34 (1956) 957, 964. 473 D . Timm, Acta Chem. Scand. 20 (1966) 2219. 474 p. N . Noble and G. C. Pimentel, / . Chem. Phys. 49 (1968) 3165; V. Bondybey, G. C. Pimentel and P. N. Noble, ibid. 55 (1971) 540; P. N. Noble, ibid. 56 (1972) 2088; but see D. E. Milligan and M. E. Jacox, ibid. 53 (1970) 2034. 475 D . G. Tuck, Progress in Inorganic Chemistry, 9 (1968) 161. 476 G . C. Pimentel and A. L. McClellan, The Hydrogen Bond, Freeman, San Francisco (1960). 477 w . C. Hamilton and J. A. Ibers, Hydrogen Bonding in Solids, Benjamin, New York (1968).

1314

CHLORINE, BROMINE, IODINE AND ASTATINE: A. J. DOWNS AND C. J. ADAMS

than HF 2 ~ as regards thermal decomposition; unlike HF 2 ~, they are not easily stabilized in the crystalline phase by simple monatomic cations. For this and other reasons, the anions derived from hydrogen chloride, bromide and iodide have attracted less attention than HF 2 ~, though recent research has gone some way towards redressing this balance475. In representing the simplest framework within which hydrogen-bonding can be studied in isolation, such anions are of particular significance. By contrast, the difficulty of distinguish­ ing between the effects of hydrogen-bonding and those of other interactions necessitates a relatively pragmatic approach to such bonding in more complicated networks. Inevitably the earliest correct identification of HXY ~ species must be a matter of debate, especially since the authors who first reported the existence of salts containing an extra molecule or more of hydrogen halide did not always formulate the materials as derivatives of the H X Y - anion. Within these limitations, it appears that the earliest report was by Dilthey4™, who described the compounds [Si(acac)3]Cl,HCl and [Si(dibenz)3]Cl,HCl (acac = acetylacetonate; dibenz = dibenzoylacetonate). Similar compounds with tetraalkylammonium, pyridinium and quinolinium cations were reported shortly afterwards by Kaufler and Kunz479, who prepared and correctly formulated derivatives of the HC12 ~, HBr2 " and HI 2 - anions, and also gave notice of chloride and bromide species of the types H2X3~ and 113X4". Ephraim established the reversible nature of the formation and dis­ sociation processes480, but few investigations of these and related compounds were other­ wise undertaken until more recent times, which have witnessed considerable research activity in this area. Preparation*15 Preparative routes leading to derivatives of the HX2 ~ ions are essentially independent of the nature of X, provided that the appropriate cation is selected. The earliest re­ searchers478»479 prepared these salts by the action of the dry, gaseous hydrogen halide on amines or substituted ammonium halides, a technique which has also been used in more recent studies. A solid tetra-alkylammonium halide generally takes up excess hydrogen halide to form products, which are often liquid at room temperature, of the type R4NHWXW+1; excess hydrogen halide can be removed by pumping, and, in the absence of moisture, the final product is usually R4NHX2. An alternative technique432'475 is to dissolve the R4NX salt in the liquid hydrogen halide and then evaporate the solution. The use of an organic solvent as a reaction medium has also been reported. Thus, treatment of tropenyl methyl ether with excess hydrogen halide in ether affords crystalline tropenium + HX 2 " (X = Cl or Br), while the reaction of gaseous hydrogen iodide with tetrabutylammonium iodide in dichloromethane gives Bu 4 NHI 2 ; with chloro-tri-^-methoxyphenylmethane or 9-chloro-9-phenylxanthin in benzene solution, hydrogen chloride gives the HC12 ~ salt of the appropriate carbonium cation. As with HF 2 ~ and related species H^Ffi +1 ~, the composition of the anions in crystalline materials depends markedly on the nature of the cation, and, to some extent, on the temperature of crystallization. Few purely inorganic salts containing HX 2 ~ ions have been reported, and even some of these are of questionable authenticity. As long ago as 1881, Berthelot reported481 that ammonium bromide and hydrogen bromide could "possibly" combine, but there is no 478 w . Dilthey, Ber. 36 (1903) 9 2 3 ; Annalen, 344 (1906) 300. 479 F. Kaufler and E. Kunz, Ber. 42 (1909) 385, 2482. 480 F. Ephraim, Ber. 47 (1914) 1828. 48i M. Berthelot, Ann. Chim. Phys. 23 (5) (1881) 98.

THE HYDROGEN HALIDES

1315

record that this observation has subsequently been put to the test. On the other hand, there is unambiguous evidence that caesium salts of the anions HC12 ", HClBr - and HC1I - are formed by the direct interaction of hydrogen chloride and the appropriate caesium halide at low temperatures482'483; high dissociation pressures are reported for CsHCl2 at room temperature483. Deuterium analogues of these and other salts have also been prepared. The precipitate formed when hydrogen chloride is bubbled through a concentrated aqueous solution of caesium chloride has been the subject of some controversy475»484. Nevertheless, recent investigations indicate that two distinct crystalline phases are produced, and threedimensional X-ray studies of single crystals of the hexagonal phase establish it as CsCl,l/3[H3O.HCl2], containing the ions Cs+, H 3 0+, Cl~ andHCl 2 - 4 8 4 ; the analogous bromide compound CsBr,l/3[H3O.HBr2] has likewise been prepared and characterized. No sign of HC12 ~ species could be detected in the Raman spectra of aqueous solutions of hydrochloric acid of various concentrations up to 8 M in lithium chloride485. Various mixed anions of the type HXY - have been obtained as crystalline salts, usually by preparative methods in which a solid is treated with hydrogen halide gas, or in which a salt R4NY is dissolved in the liquid hydrogen halide HX and crystallized as R 4 NHXY. Deuterium analogues have also been obtained in a number of cases. The following anions have been reported to date: 475 Y

X

F-

ci-

Br~

ci- , B r "M -

Br \ I ~ , C N - , N O 3 - , formate, acetate i - , CN- , formate, acetate

and there is good reason to believe that this range is capable of considerable expansion. Among the species analogous to the halogen-containing anions are the following, which have been characterized only in recent years475: H(NCS)2 -, H(CN)2 ~, H(N0 2 ) 2 ~, H(N0 3 ) 2 -, H(I0 3 ) 2 -, H(carboxylate)2-, H[M(CO) 5 ] 2 - (M = Cr, Mo or W), H(OH2)2+, H(a-picoline oxide) 2 + , and H[(^-C5H5)Fe(CO)2]2 + 465. Physical Proper ties****™ >482-484 The main interest in the study of salts of HX2 ~ and HXY - anions has centred on their physical rather than chemical properties. Some of these physical properties, determined with somewhat variable degrees of certainty, are presented in Table 34. Although not in­ cluded in the table, vibrational and nqr properties of some deuterated derivatives have also been examined. Whereas salts of the HF 2 ~ and H 2 F 3 ~ anions have been the focus of a number of struc­ tural investigations using X-ray or neutron-diffraction methods, very little definitive structural information has so far been accumulated about salts containing other halogenbearing anions of the type HX2 - or HXY ~. According to three-dimensional X-ray studies, single crystals of the compounds CsX,l/3[H 3 0+HX 2 -](X = Cl or Br)484 contain strings of HX 2 - ions parallel to the c-axis. The X · · · X distance is 3-14 ± 0-02 A in the HC12 - ion 482 j . w . Nibler and G. C. Pimentel, / . Chem. Phys. 47 C1967) 710. 483 G. C. Stirling, C. J. Ludman and T. C. Waddington, / . Chem. Phys. 52 (1970) 2730; J. A . S. Smith, F. P Temme, C. J. Ludman and T. C. Waddington,/. Chem. Soc.y Faraday Trans. / / , 69 (1973) 1477. 484 L . W. Schroeder and J. A . Ibers, Inorg. Chem. 7 (1968) 594. 485 A . G. Maki and R. West, Inorg. Chem. 2 (1963) 657.

120

NMe 4 +

9-lt 9-2f

8-6ft

Cs NMe 4 + NEt 4 + NBu n 4 +

Cs + NBu n 4 +

HC1I-

—

l-54f

—

22 -7f

—

—

122

-150 145 170

275 220 180

HClBr"

—

519 7

550 ? 7 7

823,863 740 635

7

Ξ —

—

+

—

23-9 32-8

1 NEt 4 + NEt 4 + NBu n 4 +

0-77 1-60

HFC1HFBr" HFI-

HI 2 -

—

554

7-3 12-4

27-8

125

2200 -2025

1705 1890 1570,1650 1650,1730

2710 -2900 3145

1650-1700

700 770

1420

1551

733

218

NEt 4 + NBu n 4 +

1-80

3-35 (CsBr,l/3[H 3 0 + HBr 2 -])

d

730 -1530

7 7

219

1670

1473

1575

602,660

1233

^3

667

210

199

600

v2

7

9.4 11-6 12-8

28-6

29-4

3-14d (CsCl,l/3[H 3 0 + HCl 2 -]) 3-22e (NMe 4 + HCl 2 -)

2-27 (NH4HF2)

»Ί

150

+

4-28

2-31

—

X · · · X distance in HX2" anion (Ä)

7 ?

Cs

+

! 13-7 NEt 4 12-6-14-2 NBuV 14-7 [Si(acac)3] + N(C 2 D 5 )(CD 3 ) 3 +

+

10-2

Cs +

—

-Δ5° (eu)

126 160

!

37

-ΔΗ0 (kcal mol" 1 ) (kcal mol" 1 )

K+,Rb+,Cs+, N H 4 + , NMe 4 +

Counter-ion in crystalline salt

*b

Vibrational frequencies (cm *)*

AND HXY~

NMe 4 NEt 4 + NBuV NPent n 4 +

HBr 2 -

HC1 2 - **

HF 2 "

Anion

Thermodynamics of hydrogen-bond formation1

TABLE 34. PHYSICAL PROPERTIES OF HYDROGEN DIHALIDE ANIONS HX2~

b,c a,h

b,c j a,j a,j

i

a a,h

a,g a,g a,g g

c,d

a,f a,f a c

a,c,e,f

a-d

a,b

acac = acetylacetonate. * Assignments are given on the basis of the general conclusions reached in reference b. ** 35d nqr frequencies (in MHz) for solid derivatives of the H d ^ ion (temperature in °K): CsHCl2,20-47 (294); CsCl,l/3tH30+HCll], ll-89 5 (294); NMe4HCl2,19-51 (294); NEtiHCb, 11-89 (294)*. t NIUBr + HC1. t t NR4 + H a . a D. G. Tuck, Progress in Inorganic Chemistry, 9 (1968) 161. * J. W. Nibler and G. C. Pimentel, /. Chem. Phys. 47 (1967) 710. c G. C. Stirling, C. J. Ludman and T. C. Waddington, /. Chem. Phys. 52 (1970) 2730; J. A. S. Smith, F. P. Temme, C. J. Ludman and T. C. Waddington, /. Chem. Soc,Faraday Trans. II, 69 (1973) 1477. d L. W. Schroeder and J. A. Ibers, Inorg. Chem. 7 (1968) 594. β J. S. Swanson and J. M. Williams, Inorg. Nuclear Chem. Letters, 6 (1970) 271. f J. C. Evans and G. Y.-S. Lo, /. Phys. Chem. 70 (1966) 11. * J. C. Evans and G. Y.-S. Lo, / . Phys. Chem. 71 (1967) 3942. h J. A. Salthouse and T. C. Waddington, /. Chem. Soc. A (1966) 28. 1 J. C. Evans and G. Y.-S. Lo, /. Phys. Chem. 70 (1966) 543. J J. C. Evans and G. Y.-S. Lo, /. Phys. Chem. 70 (1966) 20 k C. J. Ludman, T. C. Waddington, J. A. Salthouse, R. J. Lynch and J. A. S. Smith, Chem. Comm. (1970) 405. 1 For recent calorimetric studies of the reaction X~(solv) + HCl(solv) -> HClX"(solv) seeR. L. Benoit, M. Rinfret and R. Domain, Inorg. Chem . 11 (1972) 2603.

1318

CHLORINE, BROMINE, IODINE AND ASTATINE: A. J. DOWNS AND C. J. ADAMS

and 3-35 ± 0-02 A in the HBr2 " ion; ca. 0-47 A less than the radius sum of the appropriate X ~ ions, these values imply the formation of very strong hydrogen bonds. The symmetrical structure of the HX2 ~ ion in these salts, indicated by the presence of a mirror plane perpen­ dicular to the X · · · X axis, is supported by the nqr486 and vibrational487 spectra. By contrast, the absence of such a mirror plane implies an unsymmetrical HC12 ~ anion in crys­ talline Me 4 NHCl 2 488 ; in this case the Cl · · · Cl distance is 3-22 ±0-02 A. In general, the results of X-ray484'488, neutron-scattering483, 35C1 nqr 385 ' 486 , and vibrational spectroscopic 475 · 482 ' 483 · 487 measurements give good grounds for beiieving that HX 2 ~ anions may exist either in a centrosymmetric or in a non-centrosymmetric form, the environment being the determining factor. The anions in the salts CsHCl2, Me 4 NHCl 2 and Buw4NHCl2 are thus presumed to be unsymmetrical, those in the salts Et4NHCl2, Pr 4 NHCl 2 and Pent 4 NHCl 2 to be symmetrical. Evidently, through its influence on the optimum energy of crystal-packing, the nature of the cation imposes a major constraint on the structure of the anion. Recent investigations482»483 suggest that the HC12 ~ion in CsHCl2 is non-linear with the symmetry C2v or G, while the conclusion that the salt Me 4 NHCl 2 , like KHF 2 , has virtually no residual entropy at, or approaching, absolute zero489 implies a potential energy curve for the anion having a single minimum, and so favours C2v symmetry. For salts containing centrosymmetric HC12 ~ ions, the small shift in 35C1 nqr frequency accompanying deuteration points to a flat or nearly flat potential well486. The analysis of the vibrational spectra of HX 2 ~ and H X Y - anions 475 · 482 ' 483 has proved to be unexpectedly difficult, partly because most of the cations used to stabilize the anions themselves contribute a rich assortment of bands, and partly because the bands attributable to the anions are usually very broad at room temperature. Nevertheless, considerable clarification has been achieved by Nibler and Pimentel482, who have developed an experi­ mental technique for preparing, and obtaining the infrared spectra of, the caesium salts of HC12 -, DC12 -, HClBr -, HC1I- and DC1I ~ at 20°K; under these conditions, the width of the absorption bands is considerably reduced, even to the point where some fine structure becomes apparent. Analysis of the results, leading to assignments incorporated in Table 34, indicates frequencies for the bending mode v2 of the anions which are about one-half the previously accepted values. The high intensity in infrared absorption of the overtone 2v2 (previously assigned as i>2) is attributed to the abnormally large second derivative of the transition-dipole characteristic of hydrogen bonds, and arising from the asymmetry of the potential function482. Vibrational assignments and force constants, reported else­ w h e r e 3 ^ ^ for ions of the types HX 2 " (X = Cl, Br or I), HFY~ (Y = Cl, Br or I) and HC1Y" (Y = Br, I or N0 3 ), and for some of their deuterated derivatives, are mostly subject to modification in the light of these findings. In no case do the results lend themselves to unambiguous deductions about the structures of the ions or about the form of the poten­ tial well in which the proton moves. Salts of the HX2 ~ and HXY - anions have mostly been described as white crystalline compounds. Somewhat exceptional, therefore, are the coloured tropenium (Tr) derivatives TrHX 2 , the absorption spectra of which (in dichloromethane solution) have been related to anion-cation charge-transfer processes (X = Cl or Br) or to internal transitions of the 4

»Ö C. J. Ludman, T. C. Waddington, J. A. Salthouse, R. J. Lynch and J. A. S. Smith, Chem. Comm. (1970) 405. 487 L. W. Schroeder, / . Chem. Phys. 52 (1970) 1972, 6453. 488 J. S. Swanson and J. M. Williams, Inorg. Nuclear Chem. Letters, 6 (1970) 271. 489 s.-S. Chang and E. F. Westrum, jun., / . Chem. Phys. 36 (1962) 2571.

THE HYDROGEN HALIDES

1319

anion (X = I). However, such studies are complicated by interference from absorptions originating from the cation, and also by the tendency of the anion to dissociate in solu­ tion 4 ^: HX2-+S^S,H++2Xor HX 2 - +S v* S, HX+X- (S = solvent molecule)

Thus, in contrast with HF 2 ~, other HX2 " or HXY _ ions appear to dissociate completely in aqueous solution. The *H nmr spectrum of a solution of the HC12 " or HBr2 ~ anion in a basic solvent like dimethyl sulphoxide or acetonitrile exhibits only a single resonance line attrib­ utable to the acidic proton, indicating a rapid exchange of hydrogen between HX2 ~ and either S,H+or S,HX. In common with other thermally unstable complex halides (e.g. the polyhalides), solid derivatives of the HX 2 ~ or HXY ~ anions gain in stability as the size of the cation M + increases; according to arguments of the type invoked elsewhere (see pp. 1251-2)289.293j this behaviour can be correlated with the difference between the lattice energies (£/) of the salts MHX 2 (or MHXY) and MX. Estimates of the free energy and enthalpy (ΔΗι) changes accompanying the reaction MX(s)+HX(g) ^ MHX2(s)

have been derived either from pressure-composition isotherms over a series of temperatures or from direct calorimetric measurements of the heat output. Hence, through the thermodynamic cycle475 MX(s)+HX(g)

ΔΗΧ

^MHX2(s) -U(MHX2)-2RT

U(MX)+2RT M+(g) + X-(gHHX(g)

ΔΗ 8

►M+(g)+HX2-(g)

access has been gained to the enthalpy change AH2 of the process HX(g)+X-(g)->HX 2 -(g)

Used as an index to the strength of the H-X bond in HX 2 ~, Δ// 2 is loosely termed the "hydrogen bond energy", though it is really a net enthalpy term since the H-X distance is almost certainly greater in the HX 2 " ion than in the parent HX molecule. A sufficiently large cation sees a convergence of the lattice energies U(MX) and t/(MHX 2 ), and hence of the enthalpy changes ΔΗχ and ΔΗ2. In keeping with this, — ΔΗγ for the reaction of hydro­ gen chloride with a tetra-alkylammonium chloride follows the order Buw4N + > Et 4 N + > Me 4 N + , and graphical arguments suggest that the values obtained for AHi are essentially at a maximum for the tetrabutylammonium cation. Minimal values have thus been deduced for the "hydrogen bond energies" in HC12 ~, HBr2 ~ and HI 2 ~. The results of such experi­ ments are quoted in Table 34, together with corresponding values of AG° and Δ5°, where these have been determined. For the HF 2 ~ ion, the same approach implies a "hydrogen bond energy" not much in excess of 37 kcal, whereas, it should be noted, estimates of the lattice energies of the salts MHF 2 (M = K, Rb or Cs) give a value of 58 ± 5 kcal. Since no entirely adequate explanation of this discrepancy has yet been found, doubts about the reliability of both calculations must prevail. Reference to salts including cations like tropylium and pyridinium, as well as tetra-alkylammonium, has shown, not unexpectedly, that AH\ is influenced to some extent by factors other than the bulk of the cation.

1320

CHLORINE, BROMINE, IODINE AND ASTATINE! A. J. DOWNS AND C. J. ADAMS

A further application of calculations of this type is in the prediction or rationalization of decomposition processes affecting salts of HXY~ anions289»475. If we consider the alternative decomposition paths and construct the cycle MXls) + HY(g) AH3

M*(g) + X-(g) +

j

[MHXY(S) I AH 4

HYfg) — M+(g) 4- X ( g ) + HY(g) + e

M % ) + Y(g)

+ H(g) + X(g) -f e

I

MY(s) 4- HX(g)

M+lg) +

Y"(g) -f HX(g) — M+(g) + Y(g)+HX(g)

+ e

it is evident that AH3-AH4

=

AU+AE-AD

where AU is the difference in lattice energies of MX and MY, ΔΕ the difference in electron affinities of X and Y, and ΔΖ) the difference in the dissociation energies of HX and HY 289 . For HBrCl", Δ£ is 6 kcal and AD 16 kcal. If, therefore, Δϋ is less than 10 kcal and we ignore entropy factors, the products should be MBr and HC1. Since, even for CsCl and CsBr, Δ£/ is only 5 kcal, this condition is certain to be fulfilled in tetra-alkylammonium salts. Likewise, it can be reasoned that the decomposition of the salt R4NHCINO3 should afford R4NNO3 and HC1, as is actually observed. The determining factor in the decomposi­ tion appears to be the wide spread of bond energies in the HX molecules. By contrast, it has generally been taken for granted that the thermal decomposition of polyhalides always produces the simple halide having the smallest anion289. Chemical Properties415 Relatively little attention has been paid to the chemical reactions of salts of HX2~ and HXY~ anions; information about the reactions is therefore sparse and provides little opportunity for systematic correlation. The anions are generally unstable with respect to moisture, tending to decompose in moist air with the production of HX gas. This is a special case of the general reaction between the anions and bases B (q.v.): HX 2

+B ^ B H + + 2 X -

in which the bases B and X - are in competition for the proton. The balance of such a pro­ cess depends on the HX 2 bond strength, the base strength of B and the states of the various components. Even washing with a base as mild as acetone causes Bu n 4 NHI 2 to lose HI. There appears to have been no report of the application of HX2~ and HXY~ salts in syn­ thetic chemistry, though it is possible that they may be useful as anhydrous sources of the appropriate hydrogen halide. It has also been suggested that the anions may be significant as intermediates in certain reactions. Theoretical Aspects of the Bonding in HX2~ and HXY~ Anions415 ~477 Despite the simplicity of the HX2 ~ and HXY ~ anions as isolated hydrogen-bonded entities, a generalized, accurate treatment of the bonding is still lacking. Although numer­ ous models based on electrostatic or molecular-orbital treatments have been proposed, the qualitative predictive power of these models is strictly limited, and few quantitative pre­ dictions are possible.

THE HYDROGEN HALIDES

1321

The first theory proposed for hydrogen-bonding in various complexes was founded on the electrostatic model, which relates the strength of the hydrogen bond to the polarity of the H-X bond. The advantages and disadvantages of this model have been well summarized by Pimentel and McClellan476. Published calculations of hydrogen bond strength in HX2 ~ anions have referred only to HF2 ~; although appealingly close agreement between calculated and experimental values is found, this is probably somewhat fortuitous, and cannot be taken as proof of the model's accuracy. A simple qualitative approach which assumes that the bonds in HX2~ can be treated by an LCAO-MO description has been advanced by Pimentel475»476. The atomic orbitals involved are the hydrogen Is and the halogen ηρσ orbitals, the appropriate combinations of which yield bonding, non-bonding and antibonding molecular orbitals in a three-centre scheme formally very similar to that of Fig. 4. In the ground state of the HX2 " ion the two electron pairs originally accommodated in the X~ orbitals occupy the bonding and the non-bonding molecular orbitals, giving the equiva­ lent of two H-X bonds each of order 0-5. Alternatively, in the formalism suggested by Linnett490, the structure can be represented in terms of two one-electron bonds ojXxH°XJx

Calculations based on a more sophisticated molecular-orbital description and carried out on the HF2 ~ ion suggest, inter alia, that the hydrogen 2ρπ orbitals are quite important in the bonding scheme; the potential importance of π-bonding in HX 2 _ has likewise been demon­ strated by calculations of overlap integrals in HF 2 - and HC12~. Attempts have also been made to estimate the covalent contribution to hydrogenbonding by considering a number of idealized formulations and attempting to deduce a wave equation which represents the appropriate mixing of the individual wavefunctions in such a way as to give an accurate description of the bond character. A comprehensive review of these and other calculations on the hydrogen bond is given elsewhere491. Of particular significance, however, are thefindingsof a recent analysis of the contributions to hydrogen-bonding made by coulombic interactions, electron-exchange, polarization of lone-pair electrons, dispersion forces and charge-transfer492: in summary, these are (a) that the hydrogen atom is unique because it has no inner shells and therefore the exchange energy in a hydrogen-bond is low; (b) that the coulombic energy is the largest attractive term but is not adequately represented by the dipole-dipole approximation; (c) that delocalization effects, which may be introduced into the model by charge-transfer terms, make an impor­ tant contribution to the energy only for moderately strong hydrogen bonds, as in HX2 ~ or HXY ~; and (d) that the coulombic, exchange and charge-transfer energies are all enhanced by a low ^-character for the lone-pair electrons, the cumulative effect of all three being responsible for the general feature that orbitals low in ^-character are much better acceptors with respect to hydrogen-bonding than those rich in ^-character. 2. Addition reactions in which the H-X bond is broken. Examples of such reactions are the protonation of the Group V hydrides MH3 (M = N, P or As) and of a wide range of organic bases: e.g. ΜΗ 3 +ΗΧ->[ΜΗ 4 ] + Χ-

4

*° J. W. Linnett, The Electronic Structure of Molecules: A New Approach, Methuen, London (1964). 491 S. Bratoi, Adv. Quant. Chem. 3 (1967) 209. 492 j . N. Murrell, Chem. in Britain, 5 (1969) 107.

1322

CHLORINE, BROMINE, IODINE AND ASTATINE: A. J. DOWNS AND C. J. ADAMS

and the addition to olefins, acetylenes, epoxides and related organic compounds: e.g. \>c '

>

y

v

i HX -».-C-Ci

)

H X + HX -*- 'i ,C-X

as well as to certain multiply bonded or low-valent inorganic compounds: e.g. SO3+HCI-+CISO2OH (HBNH) 3 +3HX -> (HXB-NH 2 ) 3

(though the exact nature of this product is still open to question)493 MX2+HX -> HMX 3 (M = Si or Ge)424,494

The reaction of a gaseous base B with the gaseous hydrogen halide to form a solid salt BH+X- may be discussed in terms of the following cycle289»293: B(g) + HX(g) Dissociation of HX into ^ H+and X" AHdiss

ß(g) +

^

+

x_(g)

Combination of B and H+

ΔΗ0 I BH+

™

■ -uLCBmx-]-2RT

™

+

x-(g)

The diiferences in entropy changes for processes involving different hydrogen halides are small, and accordingly the relative stability of the salt BH+X~ is primarily a function of the standard enthalpy change ΔΗ°. For halides of the same cation, AH° = -

tfL[BH+X-]

+ Δ # ° 1 Μ + constant

At 298°K A#° d l s s = 369-8, 333-5, 323-6 and 314-3 kcal for X = F, Cl, Br and I, respec­ tively. If, therefore, the solidfluorideBH+F~ is to be more stable with respect to dissocia­ tion than the solid iodide BH +1 ~, the lattice energy of BH +F ~ must exceed that of BH + I ~ by at least 55-5 kcal. The difference in lattice energies depends, however, on the effective radius of the cation BH + , r(BH + ); according to the Kapustinskii approximation, the pre­ ceding condition requires that 512

Lr(BH + )+l-33 ~* r(BH + )+2-19j

>

55 5

'

+

or that r(BH ) < ca. 1-1 Ä. Extension of this reasoning to other pairs of halides shows that the iodide should be the salt most stable with respect to dissociation into gaseous B and HX for any cation BH + with a radius greater than ca. 1 -6 A. Setting aside complica­ tions which may arise from the precise shape of the BH + cation, or from hydrogen-bonding in one or more of the halides, we find that the predictions are almost entirely fulfilled in practice. Thus, for the NH 4 + ion with a thermochemical radius of 1 -45 A, AH° = - 3 5 · 1, -42-1, -45-0 and -43-5 kcal for X = F, Cl, Br and I, respectively, while for the larger 493 E . K. Mellon, jun., and J. J. Lagowski, Adv. Inorg. Chem. Radiochem. 5 (1963) 259; K. Niedenzu and J. W. Dawson, The Chemistry of Boron and its Compounds (ed. E. L. Muetterties), p. 377. Wiley (1967). 494 o . M. Nefedov and M. N . Manakov, Angew. Chem., Internat. Edn. 5 (1966) 1021.

1323

THE HYDROGEN HALIDES

PH 4 + ion, no fluoride of which has been isolated, the values of ΔΗ° (in kcal) are: PH4C1, — 14Ό; PH4Br, — 23· 1; PH4I, —24-3. Hence, it is not surprising that experiments in­ volving the co-condensation of arsine and a hydrogen halide at 110°K yielded infrared evidence for the formation of AsH4 + in the case of the bromide and iodide only, and sug­ gested that the latter is the more stable system495. Here, as in other aspects of the acidity of the hydrogen halides, it is the relative weakness of the H-Br and H-I bonds that is the most important single factor in determining the relative magnitudes of the overall energy changes. All the hydrogen halides add to ^)C=G^ and —C=C— units in a great variety of organic compounds, including conjugated dienes, where both 1,2- and 1,4-additions are possible496»497. Under normal conditions, the addition is presumed to take place by an electrophilic mechanism; the rate-determining step is protonation of the multiply bonded system, possibly via an initially formed π-complex, the addition being completed by subse­ quent attack of the nucleophile X". \

c=c

/

/ \

H+ slow

I

-

I

x-

H - C — ce

II

II

-H-C-C-X

II

In support of this, it is found that the ease of addition increases in the series HF < HC1 < HBr < HI, which reflects the increasing acid strength of the hydrogen halides. Studies of the stereochemistry of the process reveal nofixedpattern; whereas some reactions are stereospecific, others are not. Foreseeably, since the rate-determining step involves electrophilic attack, the reaction is assisted by electron-repelling substituents and retarded by halogens or other electron-withdrawing groups attached to the ττ-bonded carbon atoms. The orienta­ tion of the hydrogen and halogen atoms in the product is usually defined by the empirical Markownikov rule, the halogen attaching itself to the site of lower electron density; this course is determined by the relative stabilities of the intermediate carbonium ions. In solution in water or hydroxylic solvents, acid-catalysed hydration

c = c —-H-c—C-OH

/

\

II

constitutes a competing reaction. Less polar solvents encourage radical-formation, and, in the presence of peroxide catalysts, hydrogen bromide has the capacity to add to unsaturated molecules by a free-radical mechanism, leading to a reversal of the normal 495 A. Heinemann, Naturwiss. 48 ^1961) 568. 496 See for example J. Hine, Physical Organic Chemistry, 2nd edn., McGraw-Hill (1962); R. T. Morrison and R. N. Boyd, Organic Chemistry, 2nd edn., Allyn and Bacon, Boston (1966); J. March, Advanced Organic Chemistry: Reactions, Mechanisms and Structure, McGraw-Hill (1968); P. Sykes, A Guidebook to Mechanism in Organic Chemistry, 3rd edn., Longmans, London (1970); C. K. Ingold, Structure and Mechanism in Organic Chemistry, 2nd edn., Bell, London (1969). 497 B. Capon, M. J. Perkins and C. W. Rees (eds.), Organic Reaction Mechanisms, Interscience (1965-7); B. Capon and C. W. Rees (eds.), Organic Reaction Mechanisms, Interscience (1968-70).

1324

CHLORINE, BROMINE, IODINE AND ASTATINE: A. J. DOWNS AND C. J. ADAMS

orientation of the added atoms. 1

RO Radical initiator \

2

/ C= C

/

+

HX — - R O H + X

+

X — -

\ I I

3

C—C— X +

I I C—C-X

I I I I

HX—- H-C—C-X4-X·

II I I (see Section 2, p. 1167). In this case, attack is initiated by a halogen atom. The virtually complete control of orientation which can be achieved in the addition of hydrogen bromide to unsaturated molecules by introducing radicals or radical-acceptors has been turned to advantage in organic synthesis. Hydrogen bromide is unique among the hydrogen halides in this respect because the steps of the free-radical mechanism are all exothermic. With hydrogen fluoride, stage (1) is strongly endothermic, and, though with hydrogen iodide this stage is energetically favoured, the iodine atoms formed are not sufficiently reactive to promote the later stages. With hydrogen chloride, radical-addition has been observed only in a few cases, but the reaction chains are usually so short at ordinary temperatures as to make this path less attractive than the electrophilic mechanism. It is likely, but by no means certain, that addition of hydrogen chloride to the multiple bonds of CO or RCN (R = H or an organic group), aided by a chloride ion-acceptor, is an essential prelude to reactions with aromatic compounds (ArH), as in the so-called Gatterman-Koch, Gatterman or Hoesch reactions: A1C13 CO+HC1

► [HCO] + AlCl 4 -

ArH >ArCHO CuCl

ZnCl 2 RCN4- HC1

ArH ► [ R - C = N H ] + Cl"

> ArCOR hydrolysis

and with other compounds, e.g. RCN + R'OH

HCl

► [RC(OR')=NH2]+Clanhydrous hydrochloride of conditions imino ester

hydrolysis

► RC(OR ) = 0

3. Substitution reactions. Such reactions inevitably entail cleavage of the H-X bond; depending on the subsequent fate of the fragments, the hydrogen halide exercises in any given reaction at least two of the following possible functions: oxidation, reduction, protonation or halogenation. HX

-

H+

Protonating agent

4-

X" Halogenating agent Oxidizing agent

Reducing agent, e.g. metal

i Oxidation: Reaction with Metals For the reaction of a gaseous hydrogen halide with an element M to proceed M+wHX->MX n +rt/2H 2

it is necessary but not sufficient that AG/[MFJ < -65-3«, AG/°[MCln] < -22-78«,

1325

THE HYDROGEN HALIDES 1

AG/fMBrJ < —12-77« or AGf°[Mln] < +0-38« kcal mol" , thermodynamic conditions which imply that most metals should react436. In practice, the facility of reaction between metals and the hydrogen halides varies markedly with the nature of the metal and with its physical state, though, with most metals in the massive state, reaction is slow at all but elevated temperatures. As noted in Table 22, reactions such as Co+2HBr

Red heat

> CoBr2+H2

250-300°C

Pu+ 3HI > Pul3 + 3/2H2 provide expedient methods of preparing anhydrous metal halides, while the analogous process „ 35o°c Si+2HX > SiX2+H2 is probably the first stage of the reaction of a hydrogen halide with silicon to produce (mainly) HS1X3 and S1X4. The formation of a relatively volatile metal halide, through the action of a gaseous hydrogen halide, has been exploited to effect vapour-phase transport of the metal (e.g. iron or nickel) at temperatures well below those required to cause signif­ icant vaporization of the pure metal425. Aqueous hydrohalic acids attack most metals in accordance with the equation M+wH30+ -> Mft+ +wH20+/i/2H2 However, the thermodynamic feasibility of the reaction depends not only on the standard electrode potential of M, but also upon the concentration of the acid, the solubility of the halide MXn, and the stability of potential complex species. The importance of this last feature is illustrated by the observation that, whereas copper is not readily attacked by hydrochloric acid, the presence of thiourea, which complexes with the Cu + ion, causes the metal to dissolve in the 1M acid with a brisk evolution of hydrogen. The thermodynamic and kinetic readiness of reaction may also depend upon the presence of an oxidizing agent, e.g. air or an added agent such as nitric acid. Detailed studies of the dissolution of metals in hydrochloric acid have shown that the rate of reaction depends on the following variables: temperature, concentration of acid, the physical form of the metal (including the mechanical or thermal treatment it may have undergone), access of the solution species to the metal surface, the nature, distribution and concentration of impurities in the metal, and the pres­ ence of complexing, oxidizing or reducing agents in the liquid phase. Of these factors, the rate at which the reactive species in solution can diffuse to the metal surface is influenced by the viscosity of the acid solution, by the presence of an oxide film on the metal surface, and by the rate at which the acid solution is stirred or the metal sample rotated. The formation of a coherent, insoluble film of halide on the surface of the metal inevitably inhibits con­ tinued dissolution. The action of impurities in the metal or aqueous phase probably depends, at least in many cases, on the effect they have on the localized electrochemical cells set up on the surface of the metal. Metals notable for their resistance to attack by hydrochloric acid have been identified in connection with the commercial production and handling of the acid. The degree of corrosion of metals by the acid may be effectively reduced in the presence of dissolved inhibitors, e.g. phenolic compounds or quinoline. Reducing Action of the Hydrogen Halides In keeping with the standard potential of the couple 4X2/X ~, the hydrogen halides become increasingly strong reducing agents in the series HCl < HBr < HI. Illustrative of

1326

CHLORINE, BROMINE, IODINE AND ASTATINE: A. J. DOWNS AND C. J. ADAMS

the redox reactions common to all three compounds are the following418'424 -426,436: R + HX -+ R-H+X· (R = H or an organiq group) 0 2 +4HX->2X 2 +2H 2 0 2P+8HX -> 2PH4X+ 3X2 M0 2 +4HX -> MX 2 +X 2 +2H 2 0 (M = Mn or Pb) Η 2 0 2 +2Η 3 0 + +2Χ" ->X 2 +4H 2 0 Χ0 3 "+6Η 3 0 + + 5Χ" ->3X 2 +9H 2 0 UF 6 +2HX->UF 4 +2HF+X 2

Other agents with the capacity to oxidize each hydrogen halide to the corresponding molecular halogen include atomic nitrogen and oxygen, ozone, fluorine, hypochlorite, vanadate, selenate and tellurate, persulphate, permanganate, periodate and ruthenium and osmium tetroxide. On the other hand, only hydrogen bromide and iodide are oxidized by hot, concentrated sulphuric acid or by chromates or chlorine, while the power of hydro­ gen iodide as a reducing agent is indicated by its reactions with sulphur, sulphur chlorides, interhalogens, bromine, iron(III), copper(II), oxy-nitrogen compounds, phosphorus(V), arsenic(V) and antimony(V) derivatives, tetrasulphur tetranitride and organo-halogen compounds, e.g.418.424-426,436 2X- + C12

aq soln

^X2+2C1"

(X = Br or I; primary step in the manufacture of bromine and iodine—see Section 2, pp. 1136-40). 2HI+S

anhydrous Λ

aq soln

S0C1 2 +4H 3 0 + + 6IN 2 0+10H 3 O + +81N02-+2H30++I"

H3As04+2H30++2IS4N4+24H30+ +201RI+H 3 0 + +I"

I2+H2S

► H 2 S+2Cr + 5H 2 0+3I 2 ► 2NH4+ + 11H 2 0+4I 2 >NO+3H 2 0+iI 2 acid

; = ± H3As03+3H20+I2

alkali

► 4H2S+4NH4+ +24H20+10I2 »RH+H20+I2

Nitric acid is reduced by excess concentrated hydrohalic acid, but the product varies from nitrosyl chloride, N0 3 - + 3C1-+4H30+ -*N0C1+C12+6H20 believed to be the principal active agent of the mixture well known as aqua regia, to nitric oxide: N0 3 -+3I-+4H 3 0 + ->NO+3/2I2+6H20 The reaction of a hydrogen halide molecule with atomic hydrogen or an organic radical R· represents a relatively well-established propagation stage of the chain reaction between the parent halogen and either H 2 or RH; as such, it has been referred to in Section 2 (pp. 1168-9). The oxidation of the hydrogen halides by molecular oxygen is accelerated by photolytic action or by the agency of various catalysts. Detailed studies suggest that the homogeneous reaction of hydrogen bromide with oxygen proceeds by the following mech­ anism: HBr+0 2 ->HOOBr HOOBr + HBr -> 2HOBr HOBr+HBr ->H 2 0+Br 2

THE HYDROGEN HALIDES

1327

Added inert gases decrease the rate, probably by accelerating the decomposition of the intermediate HOOBr. By contrast, the pathway of the photochemical oxidation of hydro­ gen iodide is believed to be hv

Initiation

HI

->Η·+Ι·

Propagation

H · + HI

-> H2+1 ·

H+O2

->H0 2 -

H02+HI-*H202+l· Termination

I ·+1 ·

-+12

As in the hydrogen-oxygen reaction, H0 2 * plays a central role in the mechanism. Under forcing conditions, oxidation of the halide ions may give rise to oxyhalogen species: thus, depending on the exact conditions of concentration and current density, anodic oxidation of aqueous hydrochloric acid may yield either chloric or perchloric acid, while, in the presence of potassium persulphate, a mixture of hydriodic acid and silver nitrate is oxidized to the sparingly soluble periodate Ag3l0 5 . Generally, however, oxidation of halide to oxyhalogen ions by chemical means is more easily accomplished in alkaline media (see Fig. 2). Protonation and Halogenation The protonating action of aqueous hydrohalic acids is familiar through reactions such as

and

02-+2H 3 0 +

-*3H 2 0

S2- +2H 3 0 +

-> H 2 S+2H 2 0

CO32- +2H 3 0 + -* C0 2 +3H 2 0

which are significant as methods of bringing metal ions into solution (for example, in quali­ tative analysis), as wet methods of producing metal halides or halide complexes (see Table 22), and as laboratory routes to hydrogen sulphide and carbon dioxide. Anhydrous halides or oxyhalides are formed by reaction of the gaseous hydrogen halide—most com­ monly the chloride—with metal oxides at elevated temperatures436; for example, the reaction of HC1 with Sb203 to form SbCl3 is complete in 45 min at 300°C. Transport reactions have been described425 whereby a metal initially in the form of an involatile oxide, e.g. BeO, AI2O3 or Ti0 2 , is converted at elevated temperatures into a relatively volatile chloride or oxychloride by a stream of gaseous hydrogen chloride. Crystals of a material like FeOCl have thus been prepared. Nitrides, borides, suicides, germanides and certain carbides are also susceptible to protonation by the hydrogen halides in solution or in the gaseous phase, usually with the formation of the corresponding hydrides, e.g. Mg3N2+6HX -> 3MgX2+2NH3

though agents less volatile and less prone to undergo side-reactions are generally pre­ ferred for the preparation of these hydrides. In the chemistry of boranes, however, an­ hydrous hydrogen chloride is widely favoured as a means of protonating anionic derivatives

C.I.C. VOL II—TT

1328

CHLORINE, BROMINE, IODINE AND ASTATINE: A. J. DOWNS AND C. J. ADAMS

to produce neutral molecules, e.g. „Hi

HC1

B10H12(SEt2)2498c

The halogenating action of gaseous hydrogen halides is uppermost in reactions such as MH4+HX ->MH3X+H2

and

MH 3 X+ HX -> MH2X2+H 2

(M = Si or Ge but not C)

Catalysed by the appropriate aluminium halide, these provide a useful method of synthesiz­ ing halogen-substituted silanes and germanes. Likewise, the gaseous hydrogen halides halogenate diborane and other neutral boranes to give, for example, the terminally sub­ stituted derivatives B2H5X (X = Br or I)499. Probably as a sequel to protonation, metalcarbon bonds in many organometallic compounds, C-O bonds in alcohols and ethers, and C-N bonds in diazoketones and tertiary aromatic amines are also subject to halogenation by the hydrogen halides under various conditions, e.g. SnPh 4 +HX 2(7r-CH2CHCH2)2Ni+2HC1 (R 3 P) 2 PtMe 2 +HCl ROH+HX ROR+HX

aq soln

-* [(7r-CH2CHCH2)NiCl]2+2CH2 = CHCH 3 *oi

conditions

anhydrous

-► (R3P)2PtMeCl + MeH5oi

conditions

anhydrous ->RX+H20496 conditions cone aqueous acid

ArNR 2 +2HX

> Ph3SnX+PhH500 anhydrous

► ROH4-R'X

cone aqueous acid

(X = Br or I; R' = alkyl group)«**

► R X + [ArNH 2 R] + X

( X = Br ΟΓ 1)496

By these means, certain organometallic halides and organo-halogen compounds are ex­ pediently prepared, while the cleavage of methoxy groups by constant-boiling hydriodic acid forms the basis of the Zeisel method of estimating such groups in aromatic ethers. Further, reactions of this kind are probably involved in the degradation of naturally occur­ ring organic materials, e.g. cellulose, starch and gelatin, by the hydrogen halides, either in the gaseous or concentrated aqueous phase. Typical of the hydrogen- or halogen-exchange reactions involving the molecular hydrogen halides are the following: 498(a) J. Dobson and R. Schaeffer, Inorg. Chem. 7 (1968) 402; (b) J. Dobson, P. C. Keller and R. Schaeffer, Inorg. Chem. 7 (1968) 399; (c) M. D. Marshall, R. M. Hunt, G. T. Hefferan, R. M. Adams and J. M. Makhlouf, / . Amer. Chem. Soc. 89 (1967) 3361; (d) A. J. Downs, G. M. Sheldrick and J. J. Turner, Ann. Rep. Chem. Soc. 64A (1967) 234. 499 M. F. Hawthorne, The Chemistry ofBoron and its Compounds^. E. L. Muetterties), p. 223. Wiley (1967). 500 c . S. G. Phillips and R. J. P. Williams, Inorganic Chemistry, Vol. II, p. 563. Clarendon Press, Oxford (1966). 501 G. E. Coates, M. L. H. Green and K. Wade, Organometallic Compounds, 3rd edn., Vol. II, Methuen, London (1968).

1329

DETECTION AND ANALYTICAL DETERMINATION

gas phase D2+HX

- Hn+nx426 A1CI3,CS2 B10H14+6DC1 ^ = = ± B 1 0 H 8 D 6 + 6 H C 1 4 2 5 gas phase RH+DC1 ^ RD + HCH25.426 gas phase RX+HY -RV+HX5Q2 gas phase X2 + HY ^ XY+HX (used as bases for chemical lasers)436·502 gas or liquid MXn + wHY ^=±MXn-mYm + mHX502 phase

[R = organic group; M = B, Al, C, Si, Sn, P or As; X, Y = same or different halogen.] In addition, there have been numerous qualitative or quantitative accounts of halideexchange implicating complex halide or organo-halogen species in polar media502. The kinetic and thermodynamic properties of some of these reactions, together with the effects of chemical or photochemical catalysis, have been explored, notably with the aid of isotopically labelled species, in attempts to elucidate their mechanisms. The advantage taken of exchange reactions for isotopic substitution is illustrated by the reaction between decaborane(14) and deuterium chloride (whereby deuteration of specific sites of the B 10 frame­ work is achieved), by the preparation of deuterated benzenes through the action of deuterium chloride on benzene in the presence of aluminium chloride, and by some of the methods which have been employed to produce tritium chloride. Similarly, catalyzed exchange reactions involving hydrogen bromide afford a practical means of converting chloro- to corresponding bromo-alkanes345, though, in analogous situations including hydrogen iodide, reduction commonly prevails over halogen-exchange. 3.4. D E T E C T I O N

A N D A N A L Y T I C A L D E T E R M I N A T I O N OF H A L I D E S A N D H A L I D E IONS345,426,503,504

THE

HYDROGEN

Hydrogen Halides In commercial practice, the concentration of hydrochloric acid is commonly measured in terms of its specific gravity. In the laboratory, however, it is normal to assay hydrochloric and the other hydrohalic acids either volumetrically, e.g. by titration with standard base, or gravimetrically by precipitation of the silver halide. Hydrochloric acid is in common use as a primary or secondary standard in chemical analysis, a context in which various methods have been described for the preparation of solutions of accurately defined concentration. Impurities in the reagent-grade concentrated acid may include free halogens (< 10 - 4 %), sulphate and sulphite (each < 10~4%), bromide ( < 5 x l O ~ 3 % ) , ammonium ions (< 3x10-4%), arsenic (< 10-6%), iron ( < 2 x l 0 ~ 5 % ) , heavy metals (< 10" 4 %), extractable organic materials (< 5x 10~4%) and involatile matter (< 5x 10 _4 %). Tests of these specifications and of the optical transparency of the acid are described else­ where420»504; the nature of the manufacturing process determines which tests are most relevant. Hydrogen halides in the vapour phase, irrespective of concentration, are first trapped by absorption in a suitable medium, e.g. water, standard sodium hydroxide or 502 M. F. A. Dove and D. B. Sowerby, Halogen Chemistry (ed. V. Gutmann), Vol. 1, p. 41. Academic Press (1967).

503 G . W. Armstrong, H . H . Gill and R. F . Rolf, Treatise on Analytical Chemistry P. J. Elving and E. B. Sandell), Part II, Vol. 7, p. 335. Interscience (1961). 504

(ed. I. M . Kolthoff,

V. A. Stenger, Encyclopedia of Industrial Chemical Analysis (ed. F. D. Snell and L. S. Ettre), Vol. 8, p. 1. Interscience (1969); G. Oplinger, ibid. Vol. 9, p. 333. Interscience (1970).

1330

CHLORINE, BROMINE, IODINE AND ASTATINE: A. J. DOWNS AND C. J. ADAMS

sodium carbonate solution, which is subsequently analysed by acid-base titration or by treatment with silver nitrate solution. Detection and Separation of Halide Ions Classically, detection of the halide ions in aqueous solution is normally accomplished by precipitation of the sparingly soluble silver halide in acidic solution. The individual ions may be distinguished by the colour of the silver halide and by its solubility in ammoniacal solution, which decreases in the order AgCl > AgBr > Agl. In admixture with the other silver halides, silver chloride can be identified by its solubility in sodium arsenite solution; subsequent acidification of the solution again yields a positive test for chloride when silver nitrate is added. However, the most distinctive tests for individual halide ions are based on their redox properties. Thus, in dilute aqueous solution chloride ions are oxidized to chlorine only by the strongest oxidizing agents. Even with a dichromate and concentated sulphuric acid, a solid chloride yields, not the elemental halogen (as does a bromide or iodide), but chromyl chloride; this affords one of the more distinctive positive tests for chloride ions. By contrast, the presence of bromide, even in a solution rich in chloride ions, can be estab­ lished by treatment of the acidified solution with hypochlorite, permanganate or hydrogen peroxide; the bromine liberated is then characterized by its yellow-brown colour in carbon tetrachloride solution or by the coloration it produces with a dyestuif like fluorescein, fuchsin or ö-naphthoflavone (see Section 2, p. 1230); alternatively bromide is detected by oxidation to bromate by hypochlorite in hot alkaline solution. Iodides are oxidized to iodine by relatively mild agents, e.g. iron(III) or nitrite; the iodine is typically detected by its action on starch solution or by the colour of its solution in an organic solvent (see Section 2, p. 1230). Iodides, if present, must be removed before the test for bromine is carried out. This may be accomplished by oxidizing the iodide with iron(III) or nitrite and boiling the solution until the free iodine has been expelled. Mixtures of all three halide ions have been identified by selective oxidation, for example, with either persulphate or permanganate. Thus, in acetic acid solution, only iodide is oxidized by persulphate; acidification with sul­ phuric acid and the addition of some persulphate then oxidizes bromide but not chloride. Efficient separation of the halide ions in solution is achieved by the use of a strong-base anion-exchange resin, a solution of sodium nitrate typically being used as the eluent. After separation, the eluted halide ions can be estimated by normal methods (see below). The anions have likewise been separated by the techniques of paper and thin-layer chromatography. In either case, an individual halide can be identified by its relative retention time (Rf value), while a semi-quantitative measure of its concentration is gained from the intensity of the developed spot. Mixtures of volatile halides are usually fractionated by distillation or by the techniques of gas-liquid or gas-solid chromatography. Thus, one of the most rigorous procedures devised for the separation of the halide ions505 exploits the selective oxidation principle in conjunction with distillation of the parent halogen. First, iodide is oxidized with hydrogen peroxide, and then bromide with 50% nitric acid; at each stage, the free halogen is distilled from the mixture, condensed and trapped in scrubbing bottles con­ taining hydrazine sulphate solution, and then estimated as the halide either by potentiometric or by turbidimetric measurements. Estimation of the Halide Ions345,426,503,504 The principal methods applicable to the quantitative analysis of chloride, bromide and iodide ions are summarized in Table 35. In practice, the anions are usually determined 505 T. j . Murphy, W. S. Clabaugh and R. Gilchrist, / . Res. Nat. Bur. Stand. 53 (1954) 13.

(a) End-point detected using K.2Cr04 as indicator, pH 5-7 (Mohr titration); unsatisfactory for iodides. (b) End-point detected using an adsorption indicator, e.g. fluorescein, dichlorofluorescein, eosin, diphenylcarbazone or p-ethoxychrysoidinee (Fajans method). (c) End-point found by adding excess AgNC>3 and then backtitrating with a standard thiocyanate solution using iron(III) as the indicator (Volhard method); widely used to estimate total halide concentration. (d) End-point determined potentiometrically (may be used to estimate mixtures of halide ions). Differential electrolytic potentiometryf enables the end-point to be located with enhanced precision; this method has been used to determine nanogram amounts of halide at extreme dilution. (e) End-point determined by amperometric methods, e.g. the dead-stop technique. First excess of Hg2+ ions detected typically using diphenylcarbazide or diphenylcarbazone as indicator; conductimetric methods have also been used.e Basis of van der Meulen method and numerous variations of this method. Br~ oxidized to Br03~, which is then estimated by iodimetric titration. Widely used for the determination of bromide in the presence of chloride.

Silver nitrate

Mercury(II) nitrate or perchlorate Hypochlorite

Precipitation

Complex-formation

Oxidation-reduction

Volumetric

Volumetric

Volumetric

2. Cl-,Br-,I~

3. Cl-,Br",I-

4. Br-

Potassium per­ Br~ oxidized to BrCN. The reaction can be monitored potentio­ manganate, metrically or the BrCN estimated by iodimetric titration; chromic acid and to estimate traces of bromide in the presence of moderate other agents in amounts of chloride, the BrCN is distilled and determined presence of CN~ | potentiometrically by a sensitive null-point method.d

Still probably the most accurate method of determining halide ions.

Comments

Silver nitrate

Reagent

Precipitation

Reaction type

Gravimetric

Method

1. Cr,Br-,I~

Halide ions

TABLE 35. METHODS AVAILABLE FOR THE QUANTITATIVE ESTIMATION OF CHLORIDE, BROMIDE AND IODIDE IoNs a ~ d

to

Potential difference measured between reference electrode and sensing electrode which is reversible with respect to halide ions, e.g. Ag/AgX. Method easily automated and can be used to monitor halide ion concentrations, for example, in in­ dustrial wastes. Constant-current and constant-potential versions of the method have been used. Electrolytically generated reagents, e.g. Ag + or Hg22 + , have been employed as the titrant to determine halide ions in this way. Such methods have the advantage of precision, convenience and rapidity of determination, and ease of auto­ mation; applicable to low concentrations and to halide ions in admixture.

Oxidation-reduction

Electrolysis

Potentiometric

Coulometric

9. C l ~ , B r - , I -

Favoured for the determination of chloride in trace amounts.

8. C l - , B r - , I -

Silver nitrate

Precipitation

Nephelometric or turbidimetric

Free halogen (usually Br2 or I2) estimated either dissolved in an organic solvent or following reaction with a suitable reagent, e.g. rosaniline, phenol red, starch or palladium iodide (see p. 1230). Optical density due to halide complex determined at a suitable frequency. A variation of this method is involved in the estimation of Cl~ with Hg(SCN) 2 , whereby the SCN" liberated is determined colorimetrically with Fe 3 + .

7. C l - , B r - , I "

Miscellaneous

Complex-formation

Oxidation, e.g. with hypochlorite or nitrite

Oxidation-reduction

Spectrophotometric

6. C l - , B r - , I "

I~ oxidized to species such as ICh". End-point determined typically by the disappearance of free iodine from an organic solvent in contact with the reaction mixture. I" oxidized to IO3", which is then estimated by iodimetric titration.

I" oxidized to I2, which is estimated with standard thiosulphate or arsenite solution.

Potassium iodate, nitrite and other agents in acid solution Potassium iodate in a strongly acid medium Chlorine or bromine water, hypochlorite or potassium per­ manganate

Oxidation-reduction

Comments

Reagent

Reaction type

Volumetric

Method

5. I -

Halide ions

TABLE 35 {cont.)

Miscellaneous

Miscellaneous or no change

Radiochemical: (a) isotope dilution

(b) neutron activa­ tion analysis

Radioactive halogen isotopes, e.g. 131I, have been used for the low-level detection and estimation of halogens, in conjunction with precipitation, solvent-extraction or ion-exchange pro­ cedures. Hence the iodine content of the thyroid gland has been measured. Used to estimate microgram or sub-microgram quantities of halogens, e.g. bromide in water or biological material. Chemical separation is necessary if mixtures of halides are to be analysed in this way.

Spectroscopic lines suitable for analytical work occur in the ultraviolet; these are produced in emission, for example, by the "copper spark" technique. The method is simple, rapid and applicable to mixtures of halides, but is of limited precision.ab

Spectrographic

12. C I - , Β Γ - , Γ

<^> 13. Cl-,Br-,I-

Applied principally to the determination of halogens in organic materials. Suitability of the method depends on the properties of the matrix provided by the sample. Used to determine bromine in petrol and other hydrocarbons, in blood serum, urine and tissue. A recent innovation involves analysis of chlorine and bromine in aromatic hydrocarbons by absorption of monochromatic X-rays (K-capture spectroscopy).h

X-ray fluorescence or absorption

Most extensively used following quantitative oxidation of X" to Χθ3~ (X = Br or I), which is then reduced irreversibly back to X" at a dropping mercury electrode; microgram quantities of halide can thus be estimated. Anodic waves have also been used to estimate halide ions, e.g. Br~ in blood.b

11. Cl-,Br-,I-

| Electrolytic oxidation or reduction

Polarographic

10. Cl-,Br-,I-

Chromatographie

15. Cl-,Br-,I-

Miscellaneous

Reagent

(b) gas-liquid or gas-solid silica gel or celite

Distribution between Stationary phase two phases: may be: Amberlite anion(a) liquid-solid exchange resin, paper or silica gel

Catalytic action on redox reaction

Reaction type

The halide ions can be separated by ion-exchange, thin-layer or paper chromatography. Separation may be followed by quantitative analysis using AgNC>3 or by semi-quantitative evaluation of the intensities of coloured spots. Used to separate, identify and estimate volatile halides, notably organo-halogen derivatives.

The catalytic actions of iodide on the rate of reaction between Ce(IV) and As(III) and of bromide on that between MnC>4~ and I2 in acid solution0 provide very sensitive procedures for the estimation of the halide ion, e.g. 0005-0-45/ug of I" in natural waters or common salt can thus be analysed.0 ·* The progress of the reaction can be monitored spectrophotometrically. The catalytic action on thiocyanate oxidations3 and on the polarographic reduction of In(III)k has also been used to determine traces of iodide.

Comments

b

Mellows Comprehensive Treatise on Inorganic and Theoretical Chemistry, Supplement II, Part I, Longmans, Green and Co. (1956). Z. E. Jolles (ed.), Bromine and its Compounds, Benn, London (1966). c G. W. Armstrong, H. H. Gill and R. F. Rolf, Treatise on Analytical Chemistry (ed. I. M. Kolthoff, P. J. Elving and E. B. Sandell), Part II, Vol. 7, p. 335. Interscience (1961). d V. A. Stenger, Encyclopedia of Industrial Chemical Analysis (ed. F. D. Snell and L. S. Ettre), Vol. 8, p. 1. Interscience (1969); G. Opiinger, ibid. Vol. 9, p. 333 (1970). e K. N. Tandon and R. C. Mehrotra, Analyt. Chim. Acta, 27 (1962) 15. f E. Bishop and R. G. Dhaneshwar, Analyt. Chem. 36 (1964) 726. s C. L. Wilson, D. W. Wilson and C. R. N. Strouts, Comprehensive Analytical Chemistry, Vol. ILA, p. 206. Elsevier (1964). h W. Seaman, H. C. Lawrence and H. C. Craig, Analyt. Chem. 29 (1957) 1631. 1 H. V. Malmstadt and T. P. Hadjiioannou, Analyt. Chem. 35 (1963) 2157; T. P. Hadjiioannou, Analyt. Chim. Ada, 30 (1964) 488, 537. * K. B. Yatsimirsky, L. I. Budarin, N. A. Blagoveshchenskaya, R. V. Smirnova, A. P. Fedorova and V. K. Yatsimirsky, Zhur. analit. Khim. 18 (1963) 103. k A. J. Engel, J. Lawson and D. A. Aikens, Analyt. Chem. 37 (1965) 203.

a

Chronometrie

Method

14. Br-,I-

Halide ions

TABLE 35 (cont.)

DETECTION AND ANALYTICAL DETERMINATION

1335

volumetrically, though measurements of the highest precision still rely on the classical gravimetric procedures. The equivalence point of the titration of a halide with standard silver nitrate solution may be detected by the use of an indicator which forms a sparingly soluble silver salt, e.g. silver chromate, whose solubility product nevertheless exceeds those of the silver halides (the Mohr method), or by the use of a dyestuif like dichlorofluorescein, which is adsorbed by the silver halide and whose colour responds to the presence of excess silver ions in the reaction mixture. Various indirect methods of titration also exist, of which the best known (the Volhard method) involves the precipitation of the halide with a measured excess of silver ions, followed by back-titration of the excess with a standard thiocyanate solution. The very limited ionization of mercury(H) halides makes possible the titration of halide ions with a solution of an ionized mercury(II) salt, e.g. the nitrate or perchlorate; diphenylcarbazide or diphenylcarbazone is typically used as an indicator to detect the first excess of mercury(II) ions. Numerous volumetric methods have been described in which bromide is subjected to oxidation to bromine, hypobromite, bromine cyanide or bromate. Of these, one of the most widely employed is based on the so-called van der Meulen method, whereby bromide is oxidized to bromate with alkaline hypochlorite; excess hypochlorite is removed by the action of formate, and the bromate is then determined iodimetrically. Such a procedure is particularly advantageous for the analysis of bromide as a minor constituent in the pres­ ence of chloride, e.g. in natural brines and similar salt solutions, though a more rapid procedure involves oxidation by hypochlorite or other agents in acidic solution, together with spectrophotometric analysis of the bromine thereby liberated504. A method has also been devised whereby traces of bromide can be separated from moderate amounts of chloride by oxidation with chromic acid in the presence of cyanide ions; following distil­ lation from the reaction mixture, the bromine cyanide thus formed is determined potentiometrically by a sensitive null-point method504. Iodide ion is oxidized comparatively easily to iodate by chlorine, bromine, hypochlorite or bleaching powder; after removal of the excess of oxidizing agent, the iodate is deter­ mined by conventional iodimetric titration. Hence, iodide can be determined in low concentrations, e.g. in natural brines, even in the presence of chloride and bromide. Appro­ priate also to samples containing small amounts of iodide is the liberation, by the action, for example, of iron(III), nitrite or hydrogen peroxide, of free iodine, which then lends itself to spectrophotometric analysis. The titration of halide ions with silver nitrate can be accurately monitored by potentiometric measurements; in this way it is also possible to analyse mixtures of the halide ions in a single operation. In a development of the potentiometric principle, a very small current —below the polarographic diffusion limit—is passed through a pair of indicator electrodes immersed in the solution; silver ions are thus generated electrolytically, and the concen­ tration of these ions determines the potential assumed by the electrodes. This method of so-called "diiferential electrolytic potentiometry", which has the virtues of both sensitivity and precision, has been applied to the determination of nanogram amounts of halide at extreme dilution345»506. Direct potentiometric measurements using a sensing electrode which is reversible with respect to halide ions—usually silver-silver halide—have been extensively employed in the analysis of halide ions, forming a convenient basis for the con­ tinuous and automatic monitoring of the ions in solution. One of the most sensitive 506 E . Bishop and R. G. Dhaneshwar, Analyt. Chem. 36 (1964) 726.

1336

CHLORINE, BROMINE, IODINE AND ASTATINE: A. J. DOWNS AND C. J. ADAMS

procedures for evaluating bromide or iodide depends upon the catalytic action of the ions on certain redox reactions, e.g. between iodine and permanganate or between cerium(IV) and arsenic(III); in a typical experiment the progress of the redox reaction is followed spectrophotometrically. Hence, minute amounts of bromide or iodide have been analysed, for example, in blood serum, tissues or natural waters. Other physical methods which have been applied to the estimation of halide ions (see Table 35) include polarography, X-ray fluorescence and absorption, and radiochemical procedures. Prior to analysis, organically bound halogen atoms are converted to the corresponding anions by hydrolysis, oxidation or reduction; the halide ions are then determined by one of the methods of Table 35. Relating to the decomposition stage, many methods have been described, the choice often depending on the nature of the compound to be analysed. The most commonly used methods are based on: (i) hydrolysis of labile C-X bonds; (ii) reduc­ tion, e.g. with sodium metal or sodium biphenyl in a suitable solvent; (iii) wet oxidation, e.g. with fuming nitric acid (the Carius method) or chromic acid; (iv) oxidation with sodium peroxide in a sealed bomb; (v) combustion in a stream of oxygen either with (the Pregl method) or without a catalyst; (vi) combustion in an enclosed atmosphere of oxygen (the Schöniger oxygen-flask method). Either wet oxidation or alkaline incineration is favoured for the destruction of organic matter when the halogen content of biological material is to be determined; this treatment follows the precipitation or extraction of proteins in the isolation of protein-bound halogen. Many of the recent determinations of iodine in bio­ logical material have been based on the use of radioactive isotopes of the element, com­ bined with radio-counting techniques employing either sodium iodide scintillation or Geiger-Müller counters; with the 131I isotope, for example, protein-bound iodine has thus been analysed503. 3.5. B I O L O G I C A L A C T I O N OF THE H Y D R O G E N H A L I D E S A N D H A L I D E IONS345,426,507-509

The principal hazards of concentrated solutions of the hydrohalic acids are associated with the chemical burns or dermatitis arising from contact with the skin, although the action of the fumes as a respiratory irritant must not be overlooked. Gaseous hydrogen chloride is a corrosive poison, for the fumes it produces in air are actually composed of minute droplets of concentrated hydrochloric acid. It exerts a destructive action on the mucous membrane and skin. Inhalation of excessive concentrations of the gas produces a severe irritation of the upper respiratory tract. The oedema or spasm of the larynx and inflamma­ tion of the respiratory system which follow the destructive action of the acid may ultimately prove fatal. No evidence of chronic systemic effects has been found. Concentrations of 0*13-0-2% are lethal to human beings in exposures lasting only a few minutes, while the maximum concentration which can be tolerated for exposures of 60 min is in the range 0-005-0-01%. The maximum concentration in the atmosphere permissible for normal working conditions has been set at 5 ppm. Although there are few details concerning the toxicities of hydrogen bromide and hydrogen iodide, as gases or as aqueous acids, practical experience suggests that the physiological effects resemble those due to hydrogen chloride. The principal features of the biochemistry and geochemistry of the halogens concern the halide anions and organo-halogen compounds; some of these features are summarized in Table 36508. Both chlorine, as one of the more abundant elements, and iodine, as a trace 507 508 509

M. B. Jacobs, The Analytical Toxicology of Industrial Inorganic Poisons, p. 640. Interscience (1967). H. J. M. Bowen, Trace Elements in Biochemistry, Academic Press (1966). E. J. Underwood, Trace Elements in Human and Animal Nutrition, 3rd edn., Academic Press (1971).

BIOLOGICAL ACTION OF THE HYDROGEN HALIDES AND HALIDE IONS

1337

element, are now known to be essential ingredients of biochemical processes. On the other hand, there is as yet no unequivocal evidence that bromine performs any such vital function, despite the fact that all animal tissues other than the thyroid, where the position is reversed, contain 50-100 times more bromine than iodine. The halogens are taken up by animals mainly as halide anions present, in varying proportions, in foodstuffs, table salt and water. In common with other anions, the halide ions suffer diffusion in soft tissues and enter the bloodstream, whence they are removed and actively pumped into the urine by the action of the kidney. Next to urea, chlorides are in fact the most abundant constituents of the urine. The anions are also excreted by mammals via the hair and nails, which are fed by the blood­ stream, and in which relatively high concentrations of the anions are found. Chloride ions exhibit certain unique functions, their concentration being regulated by active transport. They are found mainly, but not exclusively, in the intracellular space, and make up the bulk of the anions in blood plasma. The best known active secretion of chloride in vertebrates takes place in the lining of the stomach where specialized oxyntic cells are able to secrete approximately 0-17 M hydrochloric acid, which aids the digestion of food. The oxyntic cells are able to pump bicarbonate ions into the bloodstream and hydrogen ions into the stomach; chloride ions pass from the blood to the stomach contents to maintain electrical neutrality. It is also found that the distribution of Cl~ ions between plasma and the red cells of blood is similar and intimately related to that of the HCO3- ions: as the HCO3" concentration varies causing the ions to diffuse into or out of the red cells, electro­ lytic equilibrium is maintained by a counterbalancing migration of chloride ions. Such shifts of Cl - and HCO3 - ions between the plasma and the red cells play an important role in stabilizing the pH of the blood. Likewise, changes in the concentrations of Na + and K + ions are often accompanied by alterations in the concentrations of Cl ~ and HCO3 - ions in the extracellular fluid. These compensating movements of the diffusable Cl ~ ion through cell membranes appear to be a major feature of the chloride metabolism, which is therefore closely related to the metabolism of ions such as HCO3 ~ and Na +. In common with other ions whose concentration within a cell differs from that in the medium in which the cell lives, Cl~ ions exhibit an essential electrochemical function. The most important aspect of this function appears to be the availability of the ions as a source of free energy during cell stimulation, but the stabilization of emulsions formed by highly charged colloidal particles also contained within the cell may represent a significant secondary influence of the ions. Moreover, nearly all the essential elements, both major and minor, are believed to have one or more catalytic functions in the cell. In this context, it is noteworthy that chloride ions are reported to activate the enzyme α-amylase, but, in general, little is yet known about the chloride ion as a biochemical catalyst. Organo-chlorine compounds are also found in nature. Most of them are of fungal origin, e.g. griseofulvin (CnHnO^Cl) derived from Penicillium griseofulvum, and some are powerful antibiotics. Their function is presumably that of restricting bacterial growth in the neighbourhood of fungal hyphae. Organo-chlorine compounds are also present in some red algae, and their participation in photosynthesis has been mooted. Studies with 82Br indicate that bromine is retained for only short periods in the body tissues of mammals, and that it is excreted mostly in the urine; claims that it is concentrated in the thyroid and pituitary glands have not been substantiated508. Bromide and chloride ions readily interchange to some degree in the body tissues, so that administration of bromide results in some displacement of body chloride and vice versa. Bromides exert a prolonged depressant action upon all cerebrospinal centres, with the exception of those in the medulla,

oo

^

0-5 2-2 1-7 1-2 0002 006

0 063 0077 (0044) 16,000 4-6 3-9 (5-6) 3-3 x 106

2900 3950 1890 2-5 x 109

Mammalian blood (ppm) plasma (ppm) red cells (ppm) atoms/red cell

0-43 Accumulated by mammalian thyroid, also by hair.

2800 Highest in mammalian hair and skin.

Land animals (ppm)

1-150 Accumulated by some sponges and corals and in the tunics of ascidians.

60-1000 Accumulated in scleroproteins of many sponges and by several corals and molluscs.

5000-90,000 Highest in soft coelenterates, lowest in calcareous tissues.

Marine animals (ppm)

0-42 Accumulated by Feijoa sellowiana.

30-1500 Accumulated by brown algae and some diatoms.

15 Higher in some species of Cucurbitaceae and Chenopodiaceae.

740 Highest in brown algae.

5 ί 5 Said to be enriched in soil organic jSaid to be strongly absorbed by humus, matter. j Extensive areas of soil deficiency, resulting in mammalian goitre, exist.

2-5 4 1 6-2 0-2 65

Br

2000 Much higher in maritime and salt desert plants.

4700

130 180 10 150 7-8 19,000 1-2 100 Much higher in alkaline soils near the sea and in salt deserts. A major exchangeable anion in many soils.

Cl

Land plants (ppm)

Marine plants (ppm)

Igneous rocks (ppm) Shales (ppm) Sandstones (ppm) Limestones (ppm) Fresh water (ppm) Sea water (ppm) Air (/igm"3) Soils (ppm)

Location

TABLE 36. BIOGEOCHEMISTRY OF CHLORINE, BROMINE AND IODINE 5

SO

Algae,d-f plants,b'd animalsd-g

Terrestrial organisms,d marine organ­ isms6 Relatively harmless to organisms, but prolonged administration may cause "bromism".

Plants,1* animals0

Relatively harmless to organisms.

Reviews

Toxicity

b

Calculated values in brackets. a H. J. M. Bowen, Trace Elements in Biochemistry; Academic Press (1966). W. Stiles, Encyclopedia of Plant Physiology (ed. W. Ruhland), Vol. 4, p. 558. Springer, Berlin (1958). c E. Cotlove and C. A. M. Hogben, Mineral Metabolism (ed. C. L. Comar and F. Bronner), Vol. 2B, p. 109. Academic Press (1962). d E. J. Underwood, Trace Elements in Human and Animal Nutrition, 3rd edn., Academic Press (1971). e J. Roche, M. Fontaine and J. Leloup, Comparative Biochemistry (ed. M. Flovkin and H. S. Mason), Vol. 5, p. 493. Academic Press (1963). f T. I. Shaw, Proc. Roy. Soc. B150 (1959) 356. * J. Gross, Mineral Metabolism (ed. C. L. Comar and F. Bronner), Vol. 2B, p. 221. Academic Press (1962).

Scarcely toxic, but excessive dosages lead to the condition of "iodism".

Essential to red algae, brown algae and mammals. Thyroxine and other iodinated amino acids are found in sponges, corals, ascidians and verte­ brate thyroids.

Several brominated amino acids have been isolated from marine organisms. Unconfirmed evidence for essentiality to mammals.

Essential for angiosperms and mam­ mals; has electrochemical and prob­ ably also catalytic functions. Present in some antibiotics and pigments from fungi. Also present in com­ pounds in red algae.

Functions

1340

CHLORINE, BROMINE, IODINE AND ASTATINE: A. J. DOWNS AND C. J. ADAMS

and have been used in medicine since 1835 as mild sedatives, particularly in the treatment of nervous disorders345. However, the mechanism whereby bromide depresses nervous activity has still to be elucidated. Excessive doses of bromide can produce toxic effects, manifest in the condition known as "bromism"; symptoms commonly include dermatitis, but mental and emotional disturbances are the most prominent and serious features. Notable amongst the few naturally occurring organo-bromine compounds are the pigment Tyrian purple (6,6'-dibromo-indigo) produced by some molluscs, a brominated phenol extracted from certain species of red algae, and the scleroproteins of certain corals and horny sponges which incorporate the molecules 3-bromotyrosine and 3,5-dibromotyrosine. However, the place of bromine in the economy of such invertebrates is still uncertain. The ancient Greeks are reputed to have used burnt sponges successfully but quite empirically in the treatment of human goitre426'509. Knowledge of this fact and the dis­ covery of iodine in abundance in sponges as early as 1819 led the French physician Coindet to use salts of iodine therapeutically in the treatment of goitre. The presence of iodine in the alimentary intake of mammals is now known to be essential; a deficiency causes the thyroid to become enlarged—that is, the condition characteristic of goitre—in an attempt to secure more iodine by increasing the volume of blood traversing the gland. To counteract the iodine deficiency, prevalent in the natural foodstuffs and water supplies of many "areas, it has become common practice deliberately to add iodide via the water supplies and via "iodized" table salt. However, excessive dosage of iodide may induce the condition of "iodism", typically marked by the symptoms of headache, catarrh and dermatitis ("iodine rash"). Iodine occurs in body tissues as iodide ions and as organo-iodine compounds. In the thyroid it exists as iodide ions, mono- and di-iodotyrosine, thyroxine, tri-iodothyroxine, polypeptides containing thyroxine, thyroglobulin, and probably other iodinated compounds. The iodinated amino-acids are bound with other amino-acids in peptide linkage to form thyroglobulin, the unique iodinated protein of the thyroid gland. The chief constituent of the colloid filling the follicular lumen, thyroglobulin is a glycoprotein with a molecular weight of 650,000; it constitutes the storage form of the thyroid hormone, believed to be thyroxine, and normally accounts for some 90% of the total iodine of the gland. Details of the metabolism of the thyroid function are treated elsewhere509. Iodinated amino-acids, e.g. 3-iodotyrosine and 2-iodohistidine, are also incorporated in the sclero­ proteins of certain corals and horny sponges, but the functions of these compounds and of the iodine in brown algae remain obscure.

4. DERIVATIVES OF CHLORINE, BROMINE AND IODINE IN POSITIVE OXIDATION STATES A. HALOGEN CATIONS 5 ! 0 1. I N T R O D U C T I O N

Although in some ways the choice is arbitrary, it is nevertheless difficult to decide exactly which species and which phenomena should be discussed under the heading of "halogen cations". In a review bearing this title published in 1962510, the authors restricted 5io J. Arotsky and M. C. R. Symons, Quart. Rev. Chem. Soc. 16 (1962) 282; R. J. Gillespie and M. J. Morton, ibid. 25 (1971) 553.

THE HALOGENS IN OXIDIZING ACIDIC MEDIA

1341

themselves to the species Cl + , Br + and I + ; however, most of the evidence then construed in terms of X + has since been reinterpreted in favour of X2 + , so that the circumstantial basis for such an exclusive account no longer exists. On the other hand, it would not be feasible to cite all the instances when positively-charged halogen-containing molecules have been invoked to rationalize experimental observations. It is therefore intended that this part of Section 4 shall cover the chemistry of cationic derivatives of the positive halogens generally, with particular reference to the cations X + and X 2 + (X = Cl, Br, I) and their complexes with neutral donors (e.g. water, pyridine); the species X 3 + , being but special examples of the triatomic polyhalogen cations, are here compared with the ions XY2 + (X = F, Y = Cl; X = Cl, Y = F; X = Br, Y = F, I; X = I, Y = F, Cl, Br), and accounts are also given of the pentatomic cations I 5 + and XF 4 + (X = Cl, Br, I), of XF 6 + (X = Cl, Br, I), and of the oxohalogen cations. Organohalogen cations are discussed in detail in Section 4D. While it is only recently that many of the cations described below have been properly characterized, cationic halogen species have featured more or less plausibly in the literature over many years. Earlier reviewers510 documented (but perhaps did not evade) the "obsessive desire" to crown the increasingly electropositive character of the heavier halogens with the simple iodine cation I + ; for example, the discovery that iodine and iodine monochloride formed conducting solutions in ethanol511 engendered the belief that I + was present in these solutions. Since many deductions were made from purely chemical evidence, positive halogen compounds were often and understandably confused with cationic species: the true nature of any cations actually discovered was even less well understood. In the nomenclature of this section, cations or positively-charged species are referred to as such, to distinguish them from other "positive" halogen compounds. The criteria of cation-formation are taken largely from structural and physicochemical measurements rather than from more qualita­ tive chemical details. Before proceeding to discuss the various cations, it will be useful to outline the areas of halogen chemistry wherein evidence for cationic behaviour has been found. Since the ions are powerful electrophiles, their existence as distinct species is to be expected only in media of low nucleophilicity, i.e. in strong acids. Here they may be formed either by oxidation of a halogen or by abstraction of halide ion from an interhalogen. The ionization of halogens, interhalogens, and positive-halogen derivatives in common non-aqueous solvents of higher nucleophilicity (e.g. EtOH, MeCN) has been intimated, while kinetic measurements also imply that the species X+ (aq) and Η2θΧ + may be intermediates in the halogenation of aromatic compounds. In keeping with the increasingly electropositive character of the elements, the incidence of cation-formation increases from fluorine, for which FC1+ is the only likely candidate, to iodine, which has many cationic derivatives. 2. T H E H A L O G E N S I N O X I D I Z I N G A C I D I C M E D I A

Like the neighbouring elements antimony and tellurium, iodine dissolves in oxidizing acidic media to form highly coloured polyatomic cations: the identity of the species pro­ duced depends on the oxidizing and acidic strengths of the medium, as is nicely exemplified in the iodine-sulphuric acid system. Iodine is but slightly soluble (< 10 ~4 M) in concentrated sulphuric acid, yielding a pink solution whose absorption spectrum is akin to that of iodine vapour. Oleums are both 5Π P. Waiden, Z . phys. Chem. 43 (1903) 385.

1342

CHLORINE, BROMINE, IODINE AND ASTATINE: A. J. DOWNS AND C. J. ADAMS

stronger acids and stronger oxidizing agents than the parent sulphuric acid. The solubility limit of iodine in 30% oleum is ca, 0-5 M and the red-brown solutions contain I3 + and possibly 1 5 + : concentrations of iodine in excess of 10 M may be obtained in 65% oleum, when the deep blue paramagnetic solutions contain I 2 + and S0 2 , but in disulphuric acid (45% oleum) oxidation of iodine to I2 + is incomplete. In 30% oleum I3 + is readily converted to l 2 + by iodate, persulphate or peroxide; further oxidation of I 2 + by these reagents produces iodine(III), usually in the form of solvated IO + . Concentrated sulphuric acid itself is too weakly acidic to sustain I2 + , so that oxidation of iodine with, for example, iodate produces only I3 + and IO + . That strongly electrophilic iodine species are present in the brown and blue solutions in oleum was early recognized by the respective iodination of such inert aromatic compounds as nitrobenzene and pyridine. The identity of the cations present and the equilibria relating them in oleums and other solvents, including the "magic acid" systems and aprotic media such as IF5 and SbF5, have been determined by investigation of the magnetic, colligative, conductimetric and spectroscopic properties of the solutions, as listed in Table 37: solid compounds isolated from solution also reveal the character of the halogen cations. An ele­ gant profile of the sulphuric acid systems has been presented by Gillespie and his coworkers512, who controlled the formal oxidation state of the iodine in solution by using iodine-iodic acid mixtures; in fluorosulphuric acid and related solvents S 2 0 6 F 2 is an expedient oxidant513. Cationic derivatives of chlorine and bromine are stable only in media considerably more acidic than those which support the corresponding iodine cations. Both chlorine and bromine dissolve unchanged in 65% oleum, and on further oxidation give tervalent com­ pounds, in the case of bromine via Br3 + . Br2 + and Br3 + have been characterized spectroscopically in the solvent HS0 3 F-SbF 5 -3S03 following the dissolution of bromine fluorosulphates (or Br2-S2C>6F2 mixtures)514, and are also known in the red solids [Br2]+[Sb3F16] ~ 515 and [Br3]+[AsF6] ~ 516 made from bromine, bromine pentafluoride and, respectively, antimony or arsenic pentafluoride. The order of electrophilicity Br2 + > Br3 + ~ I2 + > I3 + is indicated by the acidic strengths of the media which stabilize these species. Solutions of chlorine or chlorine fluorides in SbF5 and related solvents apparently contain two paramagnetic species in temperature-dependent equilibrium with one another; by one account517 the species predominant at -80°C is Cl 2 + and that at +60°C C1F+, although the respective identities CI2O + and ClOF + have also been proposed518. Since the only solids to have been isolated from such solutions contain diamagnetic cations like C12F + and Cl3 + , it seems probable that the paramagnetic moieties are indeed CI2O + and ClOF + , produced by reaction with the cell walls or with impurities.

512

37.

R. A. Garrett, R. J. Gillespie and J. B. Senior, Inorg, Chem, 4 (1965) 563; see also references in Table

SB F. Aubke and G. H. Cady, Inorg. Chem, 4 (1965) 269; R. J. Gillespie and J. B. Milne, ibid, 5 (1966) 1236, 1577. 514 R. J. Gillespie and M. J. Morton, Chem, Comm, (1968) 1565. 515 A. J. Edwards, G. R. Jones and R. J. C. Sills, Chem, Comm. (1968) 1527; A. J. Edwards and G. R. Jones, / . Chem, Soc, (A) (1971) 2318. 516 O. Glemser and A. Smalc, Angew, Chem,, Internat, Edn, 8 (1969) 517. 517 G. A. Olah and M. B. Comisarow, / . Amer. Chem, Soc, 91 (1969) 2172. 518 R . s. Eachus, T. P. Sleight and M. C. R. Symons, Nature, 222 (1969) 769.

I2\IO+

solvent

solvent solvent

I2

KI IC1

h+ l 2 + (with excess SbF5)

air

solvent

I2

R. A. Garrett, R. J. Gillespie and J. B. Senior, Inorg. Chem. 4 (1965) 563. J. Arotsky, H. C. Mishra and M. C. R. Symons, / . Chem. Soc. (1961) 12. R. J. Gillespie and K.C. Malhotra, Inorg. Chem. 8 (1969) 1751. R. J. Gillespie and K. C. Malhotra, Inorg. Chem. 8 (1969) 1751; R. J. Gillespie, J. B. Milne and M. J. Morton, ibid. 7 (1968) 2221. F. Aubke and G. H. Cady, Inorg. Chem. 4 (1965) 269; R. J. Gillespie and J. B. Milne, ibid. 5 (1966) 1236, 1577; R. J. Gillespie, J. B. Milne and M. J. Morton, ibid. 7 (1968) 2221. R. D. W. Kemmitt, M. Murray, V. M. McRae, R. D. Peacock, M. C. R. Symons and T. A. O'Donnell, / . Chem. Soc. (A) (1968) 862. E. E. Aynsley, N. N. Greenwood and D. H. W. Wharmby, / . Chem. Soc. (1963) 5369. R. D. W. Kemmitt, M. Murray, V. M. McRae, R. D. Peacock, M. C. R. Symons and T. A. O'Donnell, / . Chem. Soc. (A) (1968) 862; R. J. Gillespie and M. J. Morton, / . Mol. Spectro­ scopy, 30 (1969) 178.

Reference

a cond., conductivity measurements; mag., magnetic susceptibility measurements; m.w., molecular weight determinations; i.r., infrared spectroscopy; u.V., ultraviolet-visible absorption spectroscopy. b Solid contains I 2 + , Sb(lll) and Sb(V). c Solid contains I 3 + , Sb(III) and Sb(V).

I2

mag., u.V., cond., i.r. (solids), Raman.

h+

MF 5 (M = P, As,Sb,Nb,Ta)

I2

I(SbF 5 ) 2 b ; [l2] + [ S b 2 F n ] - ; ISbF5 c

mag., u.V., cond., i.r. (solids)

l3 +

solvent

I2

IF 5

cond., m.w., mag., u.V., Raman.

I2

SbF5

U.V.

cryoscopy, cond., u.V., Raman.

I3 + , I 2 + ,I4 2 + , KS0 3 F) 3

[l2] + [ M 2 F n ] (M = Sb,Ta)

[IO] + [HS 2 0 7 ]-

S206F2

I2+ I2+

h\h+ cryoscopy, cond.,

u.V., cond.

ι 2 + ,ιο +

+ i 2 cr,ici 2 +

I 2 Br I5 + , I 3 +

Techniques* cryoscopy, cond.

Solids prepared

I5M3\IO+

Species found in solution

solvent

HIO3 HIO3 HIO3 solvent HIO3

Oxidant

I2

I2 IC1 IBr I2

Substrate

HSO3F

H2S2O7 (45% oleum) 65 % oleum

30% oleum

H 2 S0 4

Solvent

TABLE 37. INVESTIGATIONS OF IODINE-CONTAINING CATIONS IN STRONG ACIDS

1344

CHLORINE, BROMINE, IODINE AND ASTATINE: A. J. DOWNS AND C. J. ADAMS

3. THE INTERHALOGENS WITH HALIDE ION-ACCEPTORS

The interhalogens interact with Lewis acids such as BF3, SbF5 or SbCls, to give crystal­ line ionic complexes: typical reactions include BrF 5 +2SbF 5 CIF3+BF3 ICl 3 +SbCl 5 IF7+AsF5

- * [BrF 4 ] + [Sb2F n ]-> [ClF 2 ] + [BF 4 r -*pCl 2 ] + [SbCl 6 ]->[IF 6 ] + [AsF 6 ]-

There is insufficient thermodynamic data about derivatives of a given Lewis acid with different interhalogens usefully to allow comparison of the relative stabilities of the different cations; for one particular cation the thermal stabilities of the complexes vary according to the acceptor strength of the Lewis acid. Some compounds have been prepared which are formally derived from unknown interhalogens such as C12F2, C13F and Br3F. Thus C1F and SbF5 afford, not the 1:1 complex ClF,SbF5 (or C1+ SbF6~), but the 2:1 adduct formulated as [Cl2F]+[SbF6]-5i9: AsF5 combines with equimolar mixtures of Cl2 and C1F to give [CI3] +[ASFÖ] " 519, and with Br2/BrF3 or Br2/BrF5 mixtures forming [Br3]+[AsF6] ~ 516. The adducts function as electrolytes in solvents such as AsF3 or anhydrous HF. X-ray and spectroscopic studies of the solids suggest that, while the best model for these compounds is an ionic lattice, there are relatively strong interactions involving halogen-bridging between the cationic and anionic units. The significant electrical conductance displayed by some liquid interhalogens (Section 4C) is consistent with self-ionization involving interhalogen cations, e.g.520 2BrF 3 ^ — B r F 2 + + B r F 4 -

4. HALOGEN CATIONS IN AROMATIC HALOGENATION

It has already been remarked that the cations I2 + , I3 + and I5 + , found in solutions of iodine in oleum, are exceedingly potent agents for iodinating even unreactive aromatic compounds. There is also evidence from kinetic studies that cationic halogen species may be involved in halogenation under milder aqueous conditions: the reader is referred to a recent review for a detailed and critical survey of this work521. It appears that the kinetic terms [HOCl][H +] and [HOBr][H + ] which are observed in halogenations effected by acidi­ fied hypochlorous and hypobromous acids represent the distinct species H 2 OCl + and H2OBr+. With the free halogen as reagent in uncatalyzed reactions the active agents are the undissociated molecules Cl2, Br2 and I2. The intervention of the aquated Cl + ion in zeroth-order chlorinations by acidified hypochlorous acid, though eloquently supported by the kinetic data, is unlikely in view of the calculated equilibrium concentration of this species (p. 1345). In many experiments silver salts have been added to limit the concentration of X", and in these cases the possibility that the complex AgX 2 + is the halogenating agent cannot be discounted^. 519 R. J. Gillespie and M. J. Morton, Inorg. Chem. 9 (1970) 811. 520 See for example A. G. Sharpe, Non-aqueous Solvent Systems (ed. T. C. Waddington), p. 285, Academic Press (1965). 521 E. Berliner, / . Chem. E4uc. 43 (1966) 124.

MONATOMIC CATIONS X +

1345

5. POSITIVE HALOGEN DERIVATIVES IN NON-AQUEOUS SOLVENTS

The chemical or electrochemical behaviour of many positive halogen derivatives dissolved in the common non-aqueous solvents implies ionization involving halogencontaining cations, although there is little justification for the belief, formerly widespread, that simple species such as I + occur in these solutions. Ethanolic solutions of INO3 may exchange up to 70% of their iodine for protons on contact with a protonated cationexchange resin, but the active agent is probably EtOHI +, the conjugate acid of ethyl hypoiodite522. When saturated solutions of I(OCOCH3)3 in acetic anhydride are electrolyzed, the amount of silver iodide formed at a silvered platinum gauze cathode is consistent with Faraday's law calculations based on P + , I3 + +Ag+3e--+AgI

The nature of the cation is still obscure. The important role of nucleophilic solvents in stabilizing halogen cations is demon­ strated by pyridine, wherein I2, IC1 and IBr dissolve with the formation of [py2I]+ cations523 : numerous derivatives of this and similar cations have been isolated and characterized (see below). The moderate electrical conductivity of solutions of IC1 in acetonitrile may similarly be explained by a partial ionization such as ICl+2MeCN ^ [ ( M e C N ) 2 I ] + + Cl-

6. MONATOMIC CATIONS X+

The cations F + , Cl + , Br + and I + have the ground state ns2np* and are highly endothermic and electrophilic species (Table 38), so that their existence as distinct chemical species would depend critically on the provision of a suitably polar, acidic environment. It was formerly supposed510 that the paramagnetic entity present in the blue solutions of iodine in strong acids was I + , but later work (Table 37) has shown that I 2 + is in fact responsible for the characteristic magnetic and spectroscopic properties of these solutions. TABLE 38. THERMODYNAMIC DATA FOR THE MONATOMIC HALOGEN CATIONS

Property

F

Cl

Br

I

A//>,o°[X+(g)](kcal mol-i)» Electron affinity of X + (kcal mol" *)» X 2 (aq) ^ X + ( a q ) + X - ( a q ) » AG°(kcalmol-i) K

422-14 402

330-74 299

301-41 273

26816 241

+ 56 10-40

+ 39 10-30

+27

IO-21

* Selected Values of Chemical Thermodynamic Properties, N.B.S. Technical Note 270-3 (1968). * J. Arotsky and M. C. R. Symons, Quart. Rev. Chem. Soc. 16 (1962) 282.

322 H . Brusset and T. Kikindai, Compt. rend. 232 (1951) 1840. 323 s . G. W. Ginn and J. L. Wood, Trans. Faraday Soc. 62 (1966) 777.

1346

CHLORINE, BROMINE, IODINE AND ASTATINE! A. J. DOWNS AND C. J. ADAMS

In media such as fluorosulphuric acid and the various oleums, iodine(I) compounds undergo, not simple ionization, e.g. IOSO2F ^ I + + S 0 3 F - ,

but complete disproportionation to I 2 + and an iodine(III) derivative. IOIOSO2F ^4l2++2I(OS02F)3 + 4S03FSimilarly equimolar mixtures of bromine and S 2 0 6 F 2 dissolved in HS0 3 F-SbF 5 -3S03 produce Br 2 + and Br(OS0 2 F) 3 rather than Br+ 519. Many kinetic studies have suggested that the aquated X + cations (X = Cl, Br, I) are intermediates in aromatic halogenations conducted in aqueous solution using the free halogen as reagent. However, the spontaneous ionic dissociation X2(aq)^X+(aq) + X-(aq) is not favoured thermodynamically (p. 1345), even allowing that the aquated X + cations, having the electronic configuration ns2np49 enjoy a ligand field stabilization of < 28 kcal m o l - i 5 i o : despite the immense uncertainty in the solvation energies assigned to the cations, the calculated equilibrium constants (Table 38) show that only I +(aq) may reason­ ably be considered as an intermediate in aqueous reactions521. The suggestion that protonated hypohalous acids H2OX + (X = Cl, Br, I) intervene in aromatic halogenations is approved on both kinetic and thermodynamic grounds510'521. Structurally there is a nice distinction between X + (aq) and H2OX + : in the former, iondipole interactions with the surrounding water molecules should leave the halogen atom with its unpaired electrons; in the latter, the localized covalent bond between H 2 0 and X h produces a diamagnetic, spin-paired species. Despite the lack of evidence for the discrete ions Cl + , Br + and I + , stable complexes of these cations with donors such as aromatic amines have been known for a long time. As was mentioned above, I2, IBr and IC1 dissolve in pyridine with the formation of the [py 2 I] + cation523. Solid derivatives are generally prepared by the action of a silver salt and the amine on the halogen in an inert solvent524. AgN03 +12+2py

CHCI3

> Agl + [py 2 I] + ΝΟ3 "

MeCN

[Ag(py) 2 ]SbF 6 + Br 2

> AgBr+[py 2 Br] + SbF6 -

AgOCOCH3 +12+py

► Agl + pyIOCOCH3

Addition of base to solutions of a halogen salt is also effective525. EtOH

C1N0 3 + 2py ► [py 2 Cl] + N0 3 ~ In proof of the presence of monovalent halogen, the iodine compounds release iodine quantitatively on reaction with aqueous iodide. LnIA+I- ->I2+A-+wL The di-ligand complexes behave as 1:1 electrolytes in solution and are best formulated as e.g. [py 2 I] + NO3 -, while compounds like pyIOCOCH3 are essentially covalent. To some extent the stoichiometry of the complex depends on the anion and it appears that the halogen cations X + have a strong tendency towards two-coordination: thus the weakly 524 H . Schmidt and H. Meinert, Angew. Chem. 71 (1959) 126. 525 M . Schmeisser and K. Brändle, Angew. Chem. 73 (1961) 388.

THE DIATOMIC CATIONS X 2 + +

1347 +

coordinating anions C104 - and SbF6 - give only [py2I] C10 4 - and [py 2 I] SbF6 -, nitrate affords both pyION0 2 and [py2I] + N 0 3 -, while with carboxylate anions the compounds are of the type [pylOCOR]. The solid "PV2I2" prepared directly from pyridine and iodine has been shown by X-ray studies to contain, in addition to I 3 _ and I 2 molecules, centrosymmetric planar [py 2 I] + cations with linear N—I—N units [r(N—I) = 2-16 Ä]526. infrared and Raman spectroscopic measurements have found essentially similar [py2Br]+ and [py2l]+ ions both in the crystal and in solution527. The {[(H2N)2CS]2I} + cation detected crystallographically in [(H2N)2CS]2l2 has a linear symmetrical S—I—S unit [r(I—S) = 2-629 A], but the ion as a whole is not planar528. 7. T H E D I A T O M I C C A T I O N S X 2 + +

Of the three cations X2 (X = Cl, Br, I), I 2 + is familiar as the origin of the paramagnetism and deep blue colour of solutions of iodine in oleum, SbF5 and similar media, where it has been defined by physicochemical studies (Table 37), including a striking resonance Raman experiment; structural details are not yet available for the solids [y+fM^Fn]"" (M = Sb, Ta)529. The red cation Br 2 + has been characterized spectroscopically in solution in the "super-acid" SbF5—HS03F—3S03514 and crystallographically in [Br2] +[Sb3F16] - 515, which is prepared from bromine, bromine pentafluoride and antimony pentafluoride. The claim that Cl2 + is found following the dissolution of chlorine or chlorine fluorides in SbF 5 and is in temperature-dependent equilibrium with C1F + is based mainly on esr measure­ ments517, and has been seriously disputed518: the solid complex Cl 2 IrF 6 has also been reported530. In the vapour phase, however, Cl 2 + has been thoroughly characterized by analysis of its electronic band spectrum. Because the ionization of the molecular halogens involves removal of an electron from an antibonding orbital, the bonds in the X 2 + cations are stronger than those in the neutral molecules: this is reflected in the higher bond dissociation energies, shorter interatomic distances and increased vibrational frequencies found in the cations (Table 39). In each case the unpaired electron produces a paramagnetic moment, measured for Br2 + and I 2 + , which is in accord with the expected 2 Π 3/2 ground state: while Cl 2 + is reputedly detectable by its esr absorption 517, the failure to record esr spectra for Br 2 + and I 2 + , even at 4°K, is attributed to a large spread in the g- and hyperfine tensors, coupled with efficient spinrelaxation530. A fully satisfactory explanation of the ultraviolet-visible spectra of Br2 + and I 2 + is still being sought. In concentrated solution I 2 + disproportionates to I 3 + and an iodine(III) species; e.g. in HSO3F the equilibrium 8I 2 + + 3S0 3 F - ^ I(S0 3 F) 3 + 5I 3 +

is satisfied by the equilibrium constant 531 [I(SQ 3 F) 3 ]*[I 3 + ]& [I 2 + ][S0 3 F-]* 526 o . Hassel a n d H . H o p e , Ada Chem. Scand. 15 (1961) 407. 527 1. H a q u e a n d J. L . W o o d , / . Mol. Structure, 2 (1968) 217. 528 H . H o p e a n d G . H.-Y. Lin, Chem. Comm. (1970) 169. 529 R . D . W. Kemmitt, M . Murray, V. M . M c R a e , R . D . Peacock, M . C . R . Symons and T. A . O'Donnell, / . Chem. Soc. (A) (1968) 862. 530 N . Bartlett, cited in ref. 529. 531 R . J. Gillespie a n d J. B . Milne, Inorg. Chem. 5 (1966) 1577.

Π3/2.«

Π, /2 ,„

Π 3 /2,Μ Πι/2,ιι

2

2

2

Anharmonic vibrational constants (cm - 1 ) 2 Π 3 /2,,

Πΐ/ 2 ,„ Dissociation energy, D0°(X2) Vibrational frequency (cm - 1 )

2

2

Χ2 + ( 2 Πι/2.«)*-Χ2θΣ ί + ) Χ 2 + (2Σ, + ) * - Χ 2 ( 1 Σ , + ) A// /t o 0 [X 2 + (g)](kcalmol-i) Dissociation energy, DQ°(X2 +) 2 Π 3 /2.*

Χ 2 + ( 2 Π 1 /2, ί ;)-Χ2( 1 Σ, + ) X 2 + ( 2 n 3 /2.«)<-X2(^ f + )

Ionization potentials: Χ 2 + ( 2 Π 3 /2,*)«-Χ2( 1 Σ, + )

Property

ω

57·1ί

C/ 2 + fe)d

35

c/ 2 + te)

β — 645-61 ω6 = 644-77 ωβ = —310

35

35-3


ωεχβ = 3-015, ajeye = 0-007

2Π3/2>5 2Π 1/2>σ 2Π„

2-479

1-53

3-98

91-8

eV

kcal

eV

14-28

362-5

15-72 265-4

321-9

13-96 253-3

32-3

76-1

kcal

329-3

286-2

ωβχΰ = 1-0 coexe = 0-35

ajexe = 1 -25

Br2 + (g)

l

L

J

| y

| S*

n

, Br 2 + AND I 2

242-4

kcal

Br 2 +

+

1-971 45-45 Br2 + (g)e Bro+ in solid ωδ [Br2][Sb3F]6V mg 376-0 368 2Π 3/2ι « 190-0 Br2 + in solutions in 2Π 1 Ρ ,„ 1520 SbF5-HS03F-3S03h\ 360

1-40

3-30

12-41

10-51

265-4

11-51

eV

kcal

eV

35C1,

TABLE 39. SOME PROPERTIES OF THE DIATOMIC CATIONS Cl2

228-3

28-8

63-9

kcal

292-0

247-7

212-6

kcal

h

(g)

—

+

1-542 35-57 I2+ in solutions in oleum, SbF* or HSÖ3F* 2IL, 238

1-25

2-77

12-66 eV

10-74

9-22

I2 +

d,e

d-h

a,b a-c

a

Reference

Πι/ 2 ,.

Π 3 /2.*

nm,g

Π3/2.„ +

d

—

018x10-6

A. W. Potts and W. C. Price, Trans. Faraday Soc. 67 (1971) 1242.

Ref.k Ref. d,e

Cl2 (g) Π 3 /2., re= 1-892 Πΐ/ 2 ,, re = 1-891 2 Π Μ re = ca. 2-25 2

2

«e

000164 000167

0-26950 0-2697 ca.019

35C12 +

in[Br2] [Sb3F16]-* 215±001

+

—

Ref. i,j

2-25 in solution* 200±01

—

—

— — —

h+

k d-f, h-j

f,i,j

d,f

d

Reference

Technical Note 270-3, U.S. Government Printing Office, Washington (1968). Data relative to Diatomic Molecules (general ed. B. Rosen), Pergamon (1970).

Ref. e,f,h

Br2 + in[Br2] + [Sb3F16]-' 1-6

Br2

+

— — —

Br 2 +

F. P. Huberman, /. Mol. Spectroscopy, 20 (1966) 29. • G. Herzberg, Molecular Spectra and Molecular Structure. I. Spectra of Diatomic Molecules, p. 520, van Nostrand (1950); Supplement to Mellor's Compre­ hensive Treatise on Inorganic and Theoretical Chemistry, Supplement II, Part I, Longmans, London (1956). f A. J. Edwards, G. R. Jones and R. J. C. Sills, Chem. Comm. (1968) 1527; A. J. Edwards and G. R. Jones, /. Chem. Soc. (A) (1971) 2318. * R. J. Gillespie and M. J. Morton, /. Mol. Spectroscopy, 30 (1969) 178. h R. J. Gillespie and M. J. Morton, Chem. Comm. (1968) 1565. 1 R. D. W. Kemmitt, M. Murray, V. M. McRae, R. D. Peacock, M. C. R. Symons and T. A. O'Donnell, /. Chem. Soc. (A) (1968) 862. 3 See references in Table 37. k G. A. Olah and M. B. Comisarow, / . Amer. Chem. Soc. 91 (1969) 2172, but see R. S. Eachus, T. P. Sleight and M. C. R. Symons, Nature, 111 (1969) 769.

b Selected Values of Chemical Thermodynamic Properties, N.B.S. c Tables of Constants and Numerical Data, No. 17, Spectroscopic d

a

Esr spectrum Ultraviolet-visible spectrum

Magnetic moment (B.M.)

Internuclear distance (Ä)

2

2

mu

2

2

Rotational constant, Be (cm - 1 )

Property

TABLE 39 (cont.)

1350

CHLORINE, BROMINE, IODINE AND ASTATINE: A. J. DOWNS AND C. J. ADAMS

At low temperatures the dimerization 2I 2 + ^ I4 2 + has been detected in HSO3F solution 5 3 2 ; the heat o f association is —10 kcal m o l - 1 . 8. TRIATOMIC CATIONS X 3 + AND XY 2 + The known triatomic halogen and interhalogen cations are listed below: the central atom is placed first. C1F2++ BrF2+ IF 2 + C12F Br3+ IC12++ CV [IBr2 ] [I2Br+] I3+ The bracketed ions have so far been found only in solution, while the others are also known in more-or-less well characterized solid derivatives, usually combined with complex halogeno-anions such as BF4 -, AsF6 ~ or SbCl6 - . CI3+ is known in the adduct [Cl 3 ] + [ A S F O ] ~ formed from Cl 2 , C1F and AsF 5 519, while [Br 3 ] + [AsF 6 ] - is prepared either from Br 2 , BrF 5 and AsF 5 or from [0 2 ] + [AsF 6 ] - and bromine 516 . Br 3 + is also known in solution in oleum 5 1 0 and in fluorosulphuric acid. First detected in oleum by its iodinating properties, I 3 + is contained in the solid I(SbF 5 ) 2 , which liberates the cation on dissolution in A s F 3 529. The black fluorosulphate I 3 O S 0 2 F is prepared from stoichiometric amounts of I 2 and S 2 0 6 F 2 5 1 3. In these cations the formally tripositive central halogen is bonded to two formally uninegative halides, giving an average oxidation state of + £ . Structural data are available for C1F 2 + , C1 2 F + , C l 3 + , BrF 2 + and IC1 2 + , which are all found to be bent, with the heaviest atom at the apex and with angles of 90-100°; the atoms of the other triatomic cations are presumably similarly disposed. X-ray investigations show a close affinity between the structures of the solids [C1F 2 ] + [SbF 6 ]-, [BrF 2 ] + [SbF 6 ] - and [IC12] + [SbCl 6 ] -. The octa­ hedral anions and angular cations are linked by bridging-halogen bonds, giving much distorted square-planar arrangements about the heavy atom (Table 40) in forming infinite helical chains (Fig. 28): in both [BrF 2 ] + [SbF 6 ] - and [IC1 2 ] + [SbCl 6 ] - the bridging utilizes cw-halogen atoms on the anion, although in [C1F2] + [SbF 6 ] ~ the anion exercises the rarer irans-bridging option, possibly because of a smaller covalent contribution to the structure. Interionic chlorine bridges are also found in [ICI2] +[A1C14] ~ 533. While the ionic model best represents the structures of these salts, the importance of interionic interactions in fixing structural details is amply demonstrated by the variations in the fundamental frequencies recorded for B r F 2 + i n a series of solids 5 3 4 and in H F - B r F 3 solution 535 . As indexed by changes in the vibraaonal frequencies of A s F 6 ~ , the extent of fluorine-bridging in the hexafluoroarsenates decreases in the order BrF 2 + > C1F 2 + > C1 2 F + > Cl 3 + 519,5359 consonant with the decreasing polarizing power of the cations. 9. PENTATOMIC CATIONS

N o structural data are available for I 5 + , the cation formed in sulphuric acid from HIO3 and an excess of iodine 512 . There has been speculation about possible shapes and electronic structures. 532 R . j . Gillespie, J. B. Milne and M. J. Morton, Inorg. Chem. 7 (1968) 2221. 533 c . G. Vonk and E. H. Wiebenga, Rec, Trav, Chim. 78 (1959) 913; Acta CrysU 12 (1959) 859. 534 K . O. Christe and C. J. Schack, Inorg, Chem, 9 (1970) 2296. 535 T . Surles, H. H. Hyman, L. A. Quarterman and A. I. Popov, Inorg, Chem. 10 (1971) 611. 536 K . O. Christe and W. Sawodny, Inorg, Chem, 8 (1969) 212; H. Meinert, U. Gross and A.-R. Grimmer, Z, Chem, 10 (1970) 226.

1351

TRIATOMIC CATIONS TABLE 40. COORDINATION ABOUT THE TRIVALENT HALOGEN IN [ClF2]+[SbF6]~

[BrF 2 ] + [SbF 6 ]- and [IC! 2 ] + [SbCl 6 ]Sb

«V' *T'

Y

a

N

b\ Sb

X

Y

a

b

(A)

(°) 95-9

92-5

Cl

F

1-58

Br

F

1-69

2-43 2-33 2-29

I

Cl

2-33 2-29

3 00 2-85

OCI

osb

OF

(a) [ClF2]+[SbF6l" projected along [100]

Reference

a

(A)

93-5

A. J. Edwards and R. J. C. Sills, / . Chem. Soc. (A) (1970) 2697. A. J. Edwards and G. R. Jones, / . Chem. Soc. (A) (1969) 1467. C. G. Vonk and E. H. Wiebenga, Acta Cryst. 12 (1959) 859.

OBr

OSb

OF

(b) [ BrFj+[SbF6 ]" projected along [010]

FIG. 28. Crystal structures of [ClF2] + [SbF6]" and [BrF 2 ] + [SbF 6 ]-

1352

CHLORINE, BROMINE, IODINE AND ASTATCNE.* A. J. DOWNS AND C. J. ADAMS

1»

0*?i!

(a)

(b) +

F I G . 29. Structures of pentatomic halogen cations: (a) possible structures of I 5 ; (b) C2t> unit of MF 4 + (M = C1, Br,I).

In so far as they can be recognized in the adducts of the halogen pentafluorides with SbF5, the ions C1F4 + , BrF4 + and IF4 + exhibit the C2v structure expected of systems in which the valence-shell of the central atom consists of four bonding electron pairs and one lone pair, and established for the isoelectronic chalcogen tetrafluorides536. Extensive fluorinebridging in the solids increases the coordination number of the halogen (Cl, 6; Br, 6; I, 8), but "axial" and "equatorial" halogen-fluorine bonds may be distinguished from the longer interionic linkages. In [IF4] +[SbF6] ~ the "axial" and "equatorial" I-F bond lengths are 1-79 A and 1-83 A, respectively, with "axial" and "equatorial" F-I-F angles of 148° and 107° 537; the crystal structure of the complex [BrF4]+[Sb2Fn] - has likewise been deter­ mined538. 10. I F 6 +

+

The discrete octahedral IFÖ cation is found in complexes of IF7 with acceptors such a& ASF5 and SbFs; 19F nmr539, vibrational540 and 129I Mössbauer541 spectroscopic studies have been reported. The molecular force field is unusual in that v2 (eg) has a higher fre­ quency than vi (aig). The cation is also formed by the oxidation of IF5 with KrF + salts. [KrFnSb 2 F n ]- +IF5 -> Kr+tIF 6 ] + [Sb2F H ]-

Analogous attempts to synthesize BrF6 + derivatives by oxidation of BrF5 have been unsuc­ cessful. 11. HALOGEN OXOCATIONS Four halogen oxocations have been reasonably well established542, viz. CIO + , CIO2 + , IO and IO2"1"; like other halogen cations, they are stable only in highly electrophilic media. They are formed (i) in the reaction of a halogen oxide with an acid or an acid anhydride, e.g. 2C102+3S03 "►[ClO][ClO2][S3O10] I203+H2Se04 -> [K>]2Se04+H20 +

or (ii) by the interaction of an oxyfluoride and a Lewis acid, e.g. SO«

[C10 2 ][S0 3 F] <

C10 2 F

AsFB

► [C10 2 ][AsF 6 ]

537 H . W. Öaird and H . F . Giles, Acta Cryst. A25 (1969) S115. 538 M . D . Lind and K . O. Christe, Inorg. Chem. 11 (1972) 608. 539 j . F . H o n and K. O. Christe, / . Chem. Phys. 52 (1970) 1960. 540 K. O. Christe, Inorg. Chem. 9 (1970) 2801. 541 S. Bukshpan, J. Soriano and J. Shamir, Chem. Phys. Letters, 4 (1969) 241. 542 M . Schmeisser and K . Brändle, Adv. Inorg. Chem. Radiochem. 5 (1963) 4 1 .

1353

THE HALOGEN-OXYGEN BOND +

There is some evidence that CIO3 may be an intermediate (i) in the preparation of CIO3F from CIO4- and acids such as HF, SbF5 or HSO3F, or (ii) in the formation of aromatic perchloryl derivatives (p. 1573) from arenes, ClC^Fand Friedel-Crafts catalysts: claims that IO3F forms unstable and presumably ionic adducts with BF 3 and AsF 5 are not unam­ biguously proven. No oxocations of bromine have been observed. In solution in HSO3F or H 2 S 2 0 7 chloryl salts function as electrolytes of strength comparable with analogous nitrosyl compounds. The vibrational spectrum of the bent discrete C10 2 + cation in the white solids [C102] +[SbF6] - and [C102]+[BF4] ~ has been reported543, although in other solids extensive cation-anion interactions occur, as witnessed by the orange colour of [C102][S03F] and the blood-red hue of [ClC^hl^Oio]. The volatility and thermal stability of the complexes [C102] +[MFn] ~ are definite functions of the acceptor strength of the parent Lewis acid: C102F,VF5 decomposes below — 78°C, while C102F,SbF5 is stable up to +230°C. The chloryl component is readily displaced by nitrogen oxides, NO

[NO] + [AsF 6 ]" + C10 2 <

[C10 2 ] + [AsF 6 ]"

NO-

y [ N 0 2 ] + [AsF6] ~ + C10 2

although the simple substitution is frequently complicated by other redox reactions. While the chlorine oxocations are relatively discrete species, IO + and IO2 + are found in solution only in solvated and associated forms: in dilute solution in oleum I2O5 is present as undissociated IO2HSO4, but at high concentrations polymerization occurs and a white solid I 2 0 5 ,S03 separates544. Similarly, conductimetric studies of I0 2 S0 3 F in HSO3F show that the solute is not completely dissociated, even at high dilution545. Spectroscopic measure­ ments suggest that solids which are formally derivatives of IO + and IO2 + [e.g. (IO)2S04, (IO)I0 3 , (I0 2 )SbF 6 , (I02)AsF6] in fact contain polymeric cationic networks with I-O-I bridges, cross-linked by the anions; in solid (IC^SC^F, however, isolated I02 + cations are joined by bridging fluorosulphate groups. B. THE OXYGEN COMPOUNDS OF THE HALOGENS 1. T H E H A L O G E N - O X Y G E N

BOND

Introduction Concern for the preservation of the octet of electrons about the halogen atom restricted early accounts of the bonding in species such as CIO4 ~ to "dative-coordinate" (1) or "semiionic" (2) electron-pair bonds. At a later stage the idea of double-bonding (3) between the halogen and oxygen atoms was mooted, but at the expense of the halogen octet. Canonical forms devised for halogen-oxygen molecules utilized all three types of bond. The applica­ tion of more sophisticated valence theory, however, prompts two important and not unconnected questions about the bonding in halogen-oxygen molecules: (1) Which orbitals on the halogen atom are involved in the bonding; specifically, is it necessary to invoke participation by halogen nd orbitals (or for iodine even 4/orbitals) in order adequately to account for observed molecular properties? (2) To what extent is π-bonding encountered in these compounds; is it p„-p„ or dn-pv in character? 543 K . O . Christe, C . J . Schack, D . Pilipovich a n d W . S a w o d n y , Inorg. Chem, 8 (1969) 2489. 544 R . j . Gillespie a n d J. B . Senior, Inorg. Chem. 3 (1964) 4 4 0 , 972. 545 H . A . C a r t e r a n d F . A u b k e , Inorg. Chem. 10 (1971) 2296.

1354

CHLORINE, BROMINE, IODINE AND ASTATINE: A. J. DOWNS AND C. J. ADAMS

Since the problem is crucial to the subsequent discussion, consideration is given at the outset to the ways in which nd orbitals may be involved in bonding; the results of experi­ ment and calculation are reviewed later.

x->o (1)

x + —o (2)

x =o (3)

It should be appreciated, however, that any discussion of this sort—and, indeed, the distinc­ tion between s, p and d electrons in many-electron systems—is an implicit acknowledgement of the one-electron approach to atomic and molecular orbitals; in this sense, many of the problems of bonding are themselves imposed by the limitations of the theoretical models currently at our disposal. nd Orbitals

The role of valence-shell d orbitals in the ground states of molecules formed by the typical elements has attracted considerable comment and attention during the past decade546 ~548. Although the nd functions undoubtedly have symmetry properties appro­ priate for bonding, there has been much debate about their potency in overlapping with ligand orbitals. The nd orbitals of the free atom are distinctly outer functions, both in size and energy. Calculations for 2nd row elements (Si, P, S, Cl) have suggested, however, that for an atom in a high valence state, or combined with highly electronegative ligands, the nd orbitals may be so stabilized and contracted that their effective participation in bonding is possible: such an effect would, in principle, lead to the involvement of nd orbitals in both σ- and ττ-interactions. The particular subjects of this section are oxyhalogen molecules, but many remarks made below are of a general nature and apply to interhalogen species and to related derivatives of the heavier elements of other Groups. The contribution of nd orbitals to σ-bonding is probably unimportant. The valencebond notion of octet expansion by means of the full hybridization of one or more nd orbitals in σ-bonding frameworks (suggested to account for coordination numbers > 4, and allowing sp*d schemes for pentacoordinate, sp3d2 schemes for hexacoordinate systems) is largely defunct, retreating before the realization that there is no causal relationship between the number of ligands and the number of bonding orbitals549. The molecular orbital approach to this problem of hypervalency550, exemplified by Rundle's three-centre-fourelectron bond551, initially ignores nd orbitals as energetically inferior to the valence-shell s and p functions; the contributions of nd orbitals and of ^-interactions (see below) are incorporated in a more sophisticated account, but may have only slight significance552 ~554. There is, however, a good a priori case for dn-pn bonding to oxygen superimposed on the molecular framework of σ-bonds546»547»555. A σ-bond between oxygen and a less electronegative halogen atom (Cl, Br, I) invariably embodies some charge separation 546 K. A . R. Mitchell, Chem. Rev. 69 (1969) 157. 547 D . W. J. Cruickshank, / . Chem. Soc. (1961) 5486. 548 L . Pauling, The Nature of the Chemical Bond, 3rd. edn., Cornell University Press, Ithaca (1960). 549 A . J. D o w n s , Unusual Coordination Numbers, in New Pathways in Inorganic Chemistry (ed. E. A . V. Ebsworth, A . G. Maddock and A . G. Sharpe), p. 15, Cambridge (1968). 550 j . I. Musher, Angew. Chem., Internat. Edn. 8 (1969) 54. 551 R. E. Rundle, Survey Progr. Chem. 1 (1963) 81. 552 L. S. Bartell, Inorg. Chem. 5 (1966) 1635. 553 B . M. D e b and C. A . Coulson, / . Chem. Soc. (A) (1971) 958. 554 R. s . Berry, M. Tamres, C. J. Ballhausen and H . Johansen, Acta Chem. Scand. 2 2 (1968) 231. 555 H. H . Jaffe, / . Phys. Chem. 58 (1954) 185.

THE HALOGEN-OXYGEN BOND

1355

X* +0* -, which is formally complete in the case of a terminal halogen-oxygen bond (2). Electron donation from a filled oxygen 2ρπ orbital to a vacant halogen nd orbital of suitable symmetry (Fig. 30) presents an attractive means of neutralizing the charge separation and presumably strengthening the bond. The extent of such interactions is expected to depend on the other groups attached to the halogen and oxygen atoms, as well as on the oxidation state and coordination number of the halogen. The number of «d orbitals able effectively to participate in dn-pn interactions may be restricted by the geometry of the molecule555. In the tetrahedral XO4 ~ anions (X = Cl, Br, I) all the nd functions possess symmetry charac­ teristics which would, in principle, allow bonding with the two filled 2pn orbitals of each

F I G . 30. Overlap for a d-κ-Ρπ bond.

oxygen atom oriented perpendicular to the X-O bond axes. As was persuasively argued by Cruickshank547, however, just the two d orbitals of E symmetry (dz2 and dxi_y2) have good overlap with the oxygen orbitals, forming two strongly bonding molecular orbitals having local 7Γ symmetry with respect to the X-O bonds; the T2 MOs are regarded as being only weakly 7r-bonding. A closely related idea, expressed in valence-bond terminology, is the duodecet rule: that the valency shell of a second row element tends to be occupied by twelve electrons556. Experimental Data The experimentally determinable parameters which bear most directly on the strength of bonding in molecules are bond lengths, valence force constants and bond energies (both dissociation energies and bond energy terms)557. Interatomic distances and valence force constants have been measured for many halogen-oxygen molecules; in Table 41 we have collated these two parameters for species containing only halogen atoms and oxygen atoms, and also for some other molecules. Bond dissociation energies, which have been less widely investigated, are listed in Table 42; bond energy terms for the chlorine oxides will be found in Table 45. The bond properties are markedly responsive to the oxidation state and coordination number of the halogen. Bonding in halogen(I) oxycompounds is effectively described by conventional electronpair bonds between the halogen and oxygen atoms. Unlike the oxygen fluorides, wherein the oxygen-fluorine bond is polarized O* +-F$ -, single bonds between oxygen and the heavier halogens are polarized Χ$+-Οδ~, a difference typified in the distinctive chemistries of 556 R. j . Gillespie and E. A . Robinson, Canad. J. Chem. 42 (1964) 2496. 557 T. L. Cottrell, The Strengths of Chemical Bonds, 2nd edn., Butterworths, London (1958).

terminal bridging terminal bridging terminal bridging

terminal bridging

terminal bridging

terminal bridging

"normal"0 h

9-20' 3-85fc 9.4u

l-40 u 1-50'

3.9v

l-70 v 1-4P 1-64*

9 0d 5.9η 9-3° 3-2° 8-2q

2.9h 3-3 1 6-41 4-3 k 7-21

1-46-1-48* 1-41° 1-71° 1-42-1-47p

1-571 1-57" 1-471

l-70

1-69

/rb

(mdyneA - 1 )

3·6 ν

l-78 r l-89 r

l-80 e l-90 e

5.9q

l-78 p l-77 r 201r 6Ό5*

1-6P

5-4* 3-0*

5-3' 3-0"

5-5n

3-91

/rb

(mdyneÄ~i)

l-82 x

l-79 m l-93 m

1-871

1-99

KL-O)* (Ä)

5-3n

2-4 3 01 4-01 4-2k

a

fr\ (mdyne A" 1 )

I

1-64-1-68*

1-721

1-82

r(Br-O)» (Ä)

Br

1 * Rounded to nearest 0 01 Ä. b Rounded to nearest 0 1 mdyne A""1. c L. Pauling, The Nature of the Chemical Bond, 3rd edn., pp. 221-230, Cornell University Press, Ithaca (1960). d K. O. Christe, C. J. Schack, D. Pilipovich and W. Sawodny, Inorg. Chem. 8 (1969) 2489. • For molecule in anhydroiodic acid HIO3J2O5; Y. D. Feikema and A. Vos, Ada Cryst. 20 (1966) 769. f N. I. Golovina, G. A. Klitskaya and L. O. Atovmyan, / . Struct. Chem. 9 (1968) 817. ' H. Siebert, Fortschr. Chem. Forsch. 8 (1967) 470. h x ~ For references see: (h) Tab?e 46; (i) Table 62; (j) Table 55; (k) Table 66; (1) Table 47; (m) Table 52; (n) Table 69; (o) Table 49; (p) Table 76; (q) Table 75; (r) Table 83; (s) Table 50; (t) Table 77; (u) Table 59; (v) Table 61; (w) Table 67; (x) Table 70.

FXO3 NXO3 2 -

(HO)5XO

HOXO3

HOX HOXO2

X0403X03X034-

X0303XOX03

xo 2 +

x2o xoxo xo 2 xo 2 02XOX02

Calculated single bond

(A)

r(Cl-O)»

Cl

TABLE 41. BOND LENGTHS AND VALENCE FORCE CONSTANTS FOR HALOGEN-OXYGEN MOLECULES

THE HALOGEN-OXYGEN BOND

1357

TABLE 42. BOND DISSOCIATION ENTHALPIES OF OXYHALOGEN COMPOUNDS

Molecule

Process

ci2o

ci2o->cio+ci

ocr ClO

oCIO-+CI+O ci-^ci+o

HOCl

cio+ ClOO OCIO C10 3

ci 2o7 HOC103

HOCl-*HO+Cl

cio+->ci++o cioo->ci+o2 ocio->cio+o cio3->ocio+o 03C10C103 -► C103+C104 HOC10 3 -^HO+C10 3

FC103 HOBr BrO

FC103-^FC102+0 HOBr-*HO+Br BrO-^Br+O

OBrO HOI IO

OBrO-> BrO+ 0 HOI-*HO+I

ιο-*ι+ο

Δ#°298°Κ

(kcalmol - 1 ) 350 32·3±2 60 ± 1 0 64-29 63-31 ±0·03(Α>°) 111 ± 5 8±2 58-7 66-5 47-6 30±4 47-6 46 57 56±10 56-23 55-2±0-6(Z) 0 °) 2s 70 56+10 43-25 42±5(ZV)

Method

Calc.a Mass spec.b Calc.c Calc.d Spectroscopice Calc.c Estimated1 Calc. a Spectroscopic* Calc. a Mass spec.b Calc.s Mass spec.h Mass spec.1 Calc. a Calc.d Spectroscopice Estimated* Calc. a Calc.d Spectroscopice

a V. I. Vedeneyev, L. V. Gurvich, V. N. Kondrat'yev, V. A. Medvedev and Ye. L. Frankevich, Bond Energies, lonization Potentials and Electron Affinities, Edward Arnold, London (1966). b I. P. Fisher, Trans. Faraday Soc. 64 (1968) 1852. c P. A. G. O'Hare and A. C. Wahl, / . Chem. Phys. 54 (1971) 3770. d Selected Values of Chemical Thermodynamic Properties, N.B.S. Technical Note 270-3 (1968). e R. A. Durie and D. A. Ramsey, Canad. J. Phys. 36 (1958) 45; R. A. Durie, F. Legay and D. A. Ramsey, ibid. 38 (1960) 444. ' S. W. Benson and J. H. Buss, / . Chem. Phys. 21 (1957) 1382. * J. B. Levy, / . Phys. Chem. 66 (1962) 1092. h G. A. Heath and J. R. Majer, Trans. Faraday Soc. 60 (1964) 1783. 1 V. H. Dibeler, R. M. Reese and D. E. Mann, / . Chem. Phys. 27 (1957) 176.

hypofluorites and hypochlorites. The three measured O-Cl(I) bond lengths (C120, 1-70 A; HOCl, 1-69 A; CH3OCl, 1-67 A) are all close to the "normal" single bond length (1-69 A) calculated from the covalent radii of the elements with due allowance for the electro­ negativity correction proposed by Schomaker and Stevenson548»558. In terms of valence force constant, however, Cl-O single bonds are less homogeneous, the values extending from 3-9 mdyne A-* (HOCl) to 2-6 mdyne A" 1 (C10S02F). Investigations of C120 using photoelectron spectroscopy and ab initio LCAO-SCF calculations imply that there is minimal (^ 1%) J-orbital participation in the C120 molecular orbitals559. The contraction of many halogen-oxygen bonds relative to the "normal" single-bond lengths and the concomitant increase in valence force constant are commonly attributed to multiple bonding546»547. The occurrence of shortening is independent of the coordination 558 v . Schomaker and D. P. Stevenson, / . Amer. Chem. Soc. 63 (1941) 37. 559 A. B. Cornford, D. C. Frost, F. G. Herring and C. A. McDowell, / . Chem. Phys. 55 (1971) 2820.

1358

CHLORINE, BROMINE, IODINE AND ASTATINE: A. J. DOWNS AND C. J. ADAMS

number of the halogen, but is restricted almost exclusively to bonds with terminal oxygen atoms; thus the bridging Cl-0 distances in C1207 are longer than the bonds in CI2O, although the terminal Cl-O distances in CI2O7 are very short. [An exception to this rule is the Cl-OH distance in anhydrous perchloric acid (1-64 A).] It is arguable to what extent the strengthening of the Cl-0 bond implied by the changes in bond length and force constant, for example in going from CIO ~ to CIO4 ~, is reflected (i) in related bond dissociation enthalpies (Table 42) or (ii) in the bond energy terms of the chlorine oxides (Table 45). Though most evidence is available for oxychlorine molecules, what data there are suggest similar diminutions in bromine-oxygen and iodine-oxygen distances as the oxidation number of the halogen rises, although the accompanying increases in stretching force constant are less than those for chlorine-oxygen bonds. The issues of interatomic distances and stretching force constants apart, evidence for nd orbital participation and multiple bonding is spectroscopic in origin, and confusing in interpretation. (i) It is claimed that the X-ray fluorescence spectra of oxyanions of second row elements (P, S and Cl) exhibit features directly testifying to 3απ-2ρπ bonding560: for oxychlorine anions the occupancy of 3d orbitals increases in the expected sequence ClO - < C102 ~ < CIO3 ~ < CIO4 ~ 561. However, X-ray spectroscopic evidence also places high positive charges on the Cl atoms of CIO3 - and CIO4 ~ (Table 44)562, implying little, if any, charge neutraliza­ tion, either by ^-donation or σ-polarization. (ii) Bond polarizabilities derived from the Raman intensities of the fundamentals in totally symmetric vibrations of the halogen oxyanions imply ηάη~2ρΉ bonding in CIO3 ~ and CIO4 -, but not in bromine and iodine systems563. (iii) The esr parameters of oxyhalogen radicals are usually interpreted with the assump­ tion of negligible spin-density in the halogen nd orbitals, although this does not necessarily exclude the participation of these functions in bonding molecular orbitals. (iv) Halogen nuclear quadrupole spectroscopy yields results which are equivocal in respect of nd orbital-participation because of the severe approximations necessary in their analysis (see pp. 1271-4). Similar reservations prevail about the interpretation of 129I Mössbauer spectra564. Calculations That 7r-bonding occurs in oxygen derivatives of intermediate and higher oxidation states of the halogens is supported, at least for oxychlorine compounds, by molecular orbital calculations of varying degrees of complexity. Wagner565 used an internally consistent LCAO-MO method to obtain ττ-bond orders and atomic charges for a number of species; the results, summarized in Table 43, show an increase in ττ-bonding in the order C l O < C120 < C102~ < ClO < CIO3- < C102 < CIO4-. No p„ character is predicted for the anions C102 ", CIO3 - and CIO4 ~, the chlorine 3p functions being pre-empted in σ-bonding orbitals and in inert pair orbitals tetrahedrally disposed about the central atom. C102 has strong d„-p„ and p„-p„ interactions, in keeping with the short interatomic distances and 560 D . S. Urch, / . Chem. Soc. (A) (1969) 3026. 56i V. I. Nefedov, / . Struct. Chem. 8 (1967) 919. 562 w . Nefedow, Phys. Status Solidi, 2 (1962) 9 0 4 . 563 G . W. Chantry and R. A . Plane, / . Chem. Phys. 3 2 (1960) 319; ibid. 3 4 (1961) 1268. 564 E . A . C . Lücken, Structure and Bonding, 6 (1969) 1. 565 E . L. Wagner, / . Chem. Phys. 37 (1962) 751.

THE HALOGEN-OXYGEN BOND

1359

TABLE 43. LCAO-MO CALCULATIONS FOR OXYCHLORINE MOLECULES*

Atomic charges Molecule

ciocl2o cio 2 cio b

C103C10 2 C104-

Qa 0 0 +0-429 +0-509 +0-293 +0-769 +0-437

Öo

rr-bond order

%3Λ character

-100 0 -0-715 -0-509 -0-431 -0-385 -0-359

0 0 0132 0-500 0-608 0-737 0-908

100 0 100 52 100

—

a

E. L. Wagner, / . Chem. Phys. 37 (1962) 751. Ab initio Hartree-Fock-Roothaan SCF-MO calculations indicate "not insignif­ icant" ^/-orbital participation; P. A. G. O'Hare and A. C. Wahl, / . Chem. Phys. 54 (1971) 3770. b

high force constant; allowing only pn-pn interactions would give a balance of one π-electron between two bonds and much weaker bonding. Detailed ab initio calculations for ClO 566 also suggest that dn-pn bonding in this molecule is "not insignificant", though Wagner predicted only ρπ-ρπ interactions. It has been shown that the inclusion of chlorine's 3d orbitals in MO calculations for the oxyanions results in lower positive charges on the Cl atoms than are assigned without the 3d functions (Table 44). While chemical intuition suggests that the halogen atom in these assemblies should not bear a charge > + 1 , X-ray spectra of the oxychlorine anions 561 · 562 are interpreted in terms of much higher charges; infrared567 and Raman 563 intensity measurements are consistent with lower charges on the Cl atoms. Comparative studies of C104 -, Br0 4 ~ and IO4 ~ [using Hartree-Fock-Slater atomic wave functions and self-consis­ tent charge and configuration (extended Hiickel) MO calculations] conclude568 that the radial nodes on the bromine Ad and iodine 5d functions do not have an adverse effect on dn-2pn overlap relative to the nodeless chlorine 3d orbitals. No attempt was made to quantify the extent of ^/-orbital participation in the bonding in perhalate anions. Bond Orders Several authors have tried to correlate the lengths r and force constants/,, of halogenoxygen bonds with their bond orders n. Pauling548 gives the following bond orders: C102~, 1-37;C10 3 -, 1-64; CIO4-, 2-10. Robinson569 discovered an empirical relation of the form logio/r(Cl-0) = -a logi 0 r(Cl-O)+6

(1)

whence he derived an equation w(Cl-O) = cft{C\-0)+d 566 p. A. G. O'Hare and A. C. Wahl, / . Chem. Phys. 54 (1971) 3770. 567 G . N. Krynauw and C. J. H. Schutte, Z. physik. Chem., N.F. 55 (1967) 121. 568 M . M. Cox and J. W. Moore, / . Phys. Chem. 74 (1970) 627. 569 E. A. Robinson, Canad. J. Chem. 41 (1963) 3021.

(2)

-0-72

a

+0-43

+0-36 a -0-26 a a

+ 100

+ 1-72

-0-43

a

ÖO

+ 102 a

+0-29

+ 0-34

ÖC1

cio 3 -

+ 2-27

+ 100 -0-82

-0-50

-0-39c -0-76d -0-78 -0-41

+ 0-57«5 + 202d + 210 + 0-65

c

b

X-ray spectroscopy

Raman intensities

LCAO-MO calculations*»6 I.r. intensities

MO calculations

LCAO-MO calculations Extended Hückel

-0-36 -0-4P

15

+ 0-44 + 0-64b

Method Estimated

Öo -0-3S

+ 0-34

Qci

C10 4 -

Not estimable because of uncertainty in "lone-pair charges". Includes chlorine 3d orbitals. Includes chlorine 4s orbitals. d Neglects chlorine 3d orbitals. β Consistent with E.S.C.A. data.

a

a

ÖO

+0-35

ÖC1

CIO 2 -

D fkfV»r p n p P I x v l W W 1 VV

R. Manne, / . Chem. Phys. 46 (1967) 4645. G. N. Krynauw and C. J. H. Schutte, Z. phys. Chem. N.F. 55 (1967) 121. G. W. Chantry and R. A. Plane, / . Chem. Phys. 34(1961) 1268. W. Nefedow, Phys. Status Solidi, 2 (1962) 904.

L. Pauling, 77te Nature of the Chemical Bond, 3rd edn., pp. 320-324, Cornell (1960). E. L. Wagner, / . Chem. Phys. 37 (1962) 751. M. M. Cox and J. W. Moore, / . Phys. Chem. 74 (1970) 627.

TABLE 44. CALCULATED AND EXPERIMENTAL ATOMIC CHARGES IN OXYCHLORINE ANIONS

THE OXIDES OF THE HALOGENS

1361

based on assumed (valence-bond) bond orders for C104" (1-5) and CIO3F (1-67) and a calculated force constant for the "normal" single bond (fr = 3-3 mdyne A -1 ). Robinson evaluated the constants: a, 645; b, 2-0; c, 0-102; and d, 0-66; using up-to-date data, the constants are: a, 6-30; b, 1 -94; c, 0-107; and d9 0-63. Wagner565 also arrived at an expression of the type (2), having calculated bond orders by the LCAO-MO method; unhappily he used a set of force constants different from those of Robinson. Siebert, too, has correlated 1-0 bond lengths and force constants570. Conclusions The deductions made about π-bonding solely on the basis of the bond lengths and force constants in halogen-oxygen compounds are inconclusive, being based on a concept of doubtful validity, viz. the "normal" single bond; while spectroscopic (X-ray, Raman etc.) measurements endorse dn-pn bonding in Cl-O systems, there is no direct physical support for ^-interactions in oxybromine or oxyiodine molecules. The variations observed in bond lengths and force constants may be attributable for I-O, for Br-O, and (at least in part) for Cl-O bonds to changes in the oxidation state and coordination number of the halogen, in the ionic charge on the assembly, and in ligation of the oxygen. In particular, the radius of the halogen atom may be very sensitive to the charge which the atom bears and to its coordination number. Scepticism may also be necessary in considering quantitative estimates of multiple-bond character. Empirically derived bond orders may be useful for comparing different halogenoxygen bonds, although this purpose might better be served by a plain statement of the physical parameters which define the bond order. Bond orders obtained in MO calculations may be more reliable, although the same calculations often produce charge distributions in the molecule very different from those determined experimentally (Table 44). In discussing oxyhalogen molecules in this section use will be made of the formulations (1), (2) and (3), since they facilitate the depiction of redox and other reactions; to do so does not imply the correctness of each or any of the formulations. 2. THE OXIDES OF THE HALOGENS INTRODUCTION

are

Fourteen halogen oxides have been isolated and reasonably well characterized: they C120

Br 2 0

CI2O3

C102 CIOCIO3

C1206 ci 2 o 7 570

Br0 2

Br 2 0 5 Br 3 0 8 Br0 3

"I2O4" I4O9 I2O5

H. Siebert, Anwendungen der Schwingungsspektroskopie in der Anorganischen Chemie* p. 106, SpringerVerlag (1966).

1362

CHLORINE, BROMINE, IODINE AND ASTATINE: A. J. DOWNS AND C. J. ADAMS

Claims have also been made for the existence of I 2 0 7 and C104, but the evidence is not compelling. It is noteworthy that the halogen perchlorates BrOC103 and 1(004)3 a r e formally mixed oxides. At room temperature the chlorine oxides are gases or fairly volatile liquids; despite their explosive nature, they have been rather fully investigated, and chlorine dioxide is used commercially as an oxidant. By contrast, the known bromine oxides are all very unstable above — 40°C, and the iodine oxides are solids which decompose into iodine and oxygen on heating. Such are the disparities (both in chemical character and in the extent of our available knowledge) between formally analogous derivatives of chlorine, bromine and iodine that the ensuing discussion is best organized in the sequence: chlorine oxides, bromine oxides, iodine oxides, rather than by comparison of compounds of the same empirical formula. CHLORINE OXIDES

Introduction

In keeping with the endothermic nature of the chlorine oxides, chlorine and oxygen do not combine under ordinary conditions: chlorine reacts with monatomic oxygen to form CI2O6, but the oxides are generally prepared by less direct chemical methods. Chlorine dioxide was discovered early in the nineteenth century during several independent investi­ gations of the action of sulphuric acid on potassium chlorate. In 1834 the oxidation of chlorine with mercuric oxide (yielding C120) was reported, while in 1843 Millon produced by the photolysis of C102 a red oil—presumably chlorine trioxide but analyzed then as ClöOn. The isolation of C1207 awaited the dehydration of perchloric acid in 1900. Two other oxides have only recently been characterized—CI2O3 in 1967 and CIOCIO3 in 1970; the unknown compound chlorine tetroxide has a colourful record of alleged preparations. The dangers inherent in the manipulation of chlorine oxides have been recognized since the earliest days of their chemistry: in 1815 Davy cautioned that the reaction of KCIO3 with sulphuric acid be performed with only very small quantities of chemicals. The need to employ adequate safety procedures when working with these unpredictable endothermic materials cannot too often be stressed. All the compounds, but especially CIO2, are liable to detonate in the event of thermal or physical shock, or even change of phase. They are strong oxidizing agents, causing organic matter to ignite spontaneously and sometimes explosively. The reader is referred to general texts on laboratory safety571, and should be fully conversant with the literature before attempting to prepare or use any of these materials. Physical and thermochemical data for the oxides of chlorine are itemized in Table 45. The account of their chemistry presented here summarizes and supplements the most recent reviews572 _575 , which together provide an exhaustive coverage of the preparations, 571

Handbook of Laboratory Safety, 2nd edn. (ed. N . V. Steere), Chemical Rubber Co., Cleveland (1971). Gmelins Handbuch der Anorganischen Chemie, 8 Auflage, "Chlor", System-nummer 6, Teil B, Liefer­ ung 2, Verlag Chemie, Weinheim/Bergstr. (1969). 573 C. C. Addison, Mellows Comprehensive Treatise on Inorganic and Theoretical Chemistry, Supplement II, Part I, pp. 514-544, Longmans, London (1956). 574 H. L. Robson, Kirk-Othmer Encyclopedia of Chemical Technology, 2nd edn., Vol. 5, pp. 7-50, Interscience, New York (1964). 575 (a) M. Schmeisser and K. Brändle, Adv. Inorg. Chem. Radiochem. 5 (1963) 4 1 ; (b) R. J. Brisdon, MTP International Review of Science: Inorganic Chemistry Series One, Vol. 3 (ed. V. Gutmann), p. 215, Butterworths and University Park Press (1972). 572

9.7f

25-09° 24-5° 28-8° 61-36° 1003° 62° 256* 79-l h

19-71* 19-2° 23-4C 63-60° 10-85°

49c

254 d

6-29f 22-23f

7-7427-1275-l/r f (227°K
red red yellow-green -59θ

6-20*> 22-5*

brown red-brown yellow-green -120-6* 2-0 b 7-87-1373/r b (173°Κ<Γ<288°Κ)

cio2

78-21 20-561

1

ca. 43 1

7-171 22-61

pale yellow pale yellow pale yellow -117±2i 44.51 7-82-1586/Γ 1 (226°Κ<Γ<296°Κ)

CIOCIO3

[ 9 1 . 4 ]l,m

[57-2]0·1 [270p.m

[37]··1

yellow (-180°C) [dark red] [red-brown] 3-5 k 203 k liq: 7-1-2070/7* solid: 9-3-2690/Γ (233°K<7 , <293°K) 9.5k 21 k

ci2o6

> 300*

65 0°

8-29n 23-4D

colourless colourless colourless -91-5n 81 n 8·03-1818/Γ η (270°K
CI2O7

C. H. Secoy and G. H. Cady, /. Amer. Chem. Soc. 62 (1940) 1036. C. F. Goodeve, / . Chem. Soc. (1930) 2733. ° Selected Values of Chemical Thermodynamic Properties, N.B.S. Technical Note 270-3 (1968). d Measured by photoelectron spectroscopy: A. B. Comford, D. C. Frost, F. G. Herring and C. A. McDowell, / . Chem. Phys. 55 (1971) 2820. β F. E. King and J. R. Partington, / . Chem. Soc. (1926) 925. f H. Grubitsch and E. Suppan, Monatsh. Chem. 93 (1962) 246. 8 Measured by mass spectroscopy: V. H. Dibeler, R. M. Reese and D. E. Mann, J. Chem. Phys. 27 (1957) 176. h Estimated from the heat of hydration of C102~: V. I. Vedeneyev, L. V. Gurvich, V. N. Kondrat'yev, V. A. Medvedev and Ye. L. Frankevich, Bond Energies, Ionization Potentials and Electron Affinities, p. 196, Edward Arnold, London (1966). 1 C. J. Schack and D. Pilipovich, Inorg. Chem. 9 (1970) 1387. 3 K. O. Christe, C. J. Schack and E. C. Curtis, Inorg. Chem. 10 (1971) 1589. k C. F. Goodeve and F. D. Richardson, / . Chem. Soc. (1937) 294. 1 Bracketed values refer to CIO3 molecules. m Estimated from the lattice energies of chlorates: ref. h. n C. F. Goodeve and J. Powney, / . Chem. Soc. (1932) 2078.

a b

Δ/fvap at b.p. (kcal mol"1) Trouton's constant (cal deg"1 mol" 1 ) Thermodynamic properties of the gaseous molecule Atf/atO^ikcalmol-i) AHf° at 298°K (kcal mol'*) AG / °at298°K(kcalmol" 1 ) 5° at 298°K (cal deg"1 mol"1) Cp° at 298°K (cal deg"1 mol" 1 ) Thermochemical properties Bond energy term (kcal mol" 1 ) Ionization potential (kcal mol" 1 ) Electron affinity (kcal mol"1)

Physical properties Colour: solid liquid vapour Melting point (°C) Boiling point (°C) Vapour pressure, log/>(mm) =

ci2o

TABLE 45. PHYSICAL PROPERTIES OF THE OXIDES OF CHLORINE

1364

CHLORINE, BROMINE, IODINE AND ASTATINE! A. J. DOWNS AND C. J. ADAMS

properties and uses of these compounds. Further historical information is best sought in the original edition of Mellor576. Dichlorine Monoxide Preparation The name "dichlorine monoxide" is adopted for the compound C120; the use of "chlorine monoxide", common in the older literature, invites confusion with the diatomic radical ClO. In the laboratory CI2O is best obtained by treating freshly prepared yellow mercuric oxide either with chlorine gas (diluted with dry air) or with a solution of chlorine in carbon tetrachloride577. 2Cl2 + 2HgO -> HgCl 2 .HgO+Cl 2 0

Large-scale preparations may also follow this method. Dichlorine monoxide is very soluble in water and is partially converted to hypochlorous acid, of which it is the anhydride. CI2O+H2O ^2HOCl

Concentrated solutions of hypochlorous acid contain significant amounts of dichlorine monoxide, and the vapour over such solutions also contains CI2O: the oxide may be extracted either with CC14 or by passing air through the solution. On dissolution in alkali, CI2O generates hypochlorite. Much of the dichlorine monoxide produced industrially is used directly in the manufacture of hypochlorous acid and hypochlorites. Structure and reactions The compound is a yellowish-green gas at room temperature. The angular Cl-O-Cl molecule is well defined in the vapour phase by spectroscopic and difTractometric techniques (Table 46); the Cl-O distance is consistent with a normal single bond. Comparison of the infrared spectra of the vapour and solid suggests that a similar C2v unit prevails in the condensed phase. C120 is liable to explode on heating or on impact, although very pure samples are allegedly quite stable in this respect. The photolytic decomposition of the gas into Cl2 and O2 proceeds with a quantum yield of 2 by the mechanism C120

^ C 1 0 + C1

CI+CI2O ->C10+C1 2

cio+cio-^ci 2 +o 2 Photolysis of carbon tetrachloride solutions or controlled thermal degradation (60°C < T < 140°C) produces significant amounts of C102. Photolysis of C120 isolated at 14°K in an inert gas matrixes affords, in addition to (C10)2, the radical CICIO, characterized by its infrared spectrum as containing a very weak Cl-Cl bond and a Cl-O linkage not much different from that in the diatomic ClO molecule (see Table 55). 576 J. W. Mellor, A Comprehensive Treatise on Inorganic and Theoretical Chemistry, Vol. 2, pp. 240-418, Longmans, Green and Co., London (1922). 577 G. H. Cady, Inorganic Syntheses, Vol. 5 (ed. T. Moeller), p. 156, McGraw-Hill (1957). 578 M. M. Rochkind and G. C. Pimentel,/. Chem. Phys. 46 (1967) 4481; W. G. Alcock and G. C. Pimentel, ibid. 48 (1968) 2373.

THE OXIDES OF THE HALOGENS

1365

TABLE 46. SPECTROSCOPIC AND RELATED INVESTIGATIONS OF DICHLORINE MONOXIDE

Microwave spectrum** Measured for 35C116()35C1, 35Cli6()37Cl, 37C116037C1 C2v symmetry ro(Cl-O) = 1-70038 ±000069 Ä; 0Oi-o-ci = Π0·96±0·08°* r s (Cl-0) = 1-70038 ± 0 0 0 0 4 3 A; 0Ci-o-ci = 110-86±0-04° a a Force constants (mdyne Ä ~ i ) : fr = 2-88±0-08; f„ = 0-31 ± 0 0 5 fy/ri = 0-432±0002; fjr = 0 1 7 ± 0 0 5 Quadrupole coupling data:* e2Qq = —140 MHz; η = 0 06 Electron diffraction0 C2v symmetry 0 C i-o-ci = 1Π·2±0·3° r,(l)(Cl-0) = l-693±0-003 Ä; Vibrational spectra* Measured for vapour, liquid, solid and matrix-isolated molecule; includes isotopic data ( 1 6 0, 1 8 0, 3 5 C1, 3 7 CI). Fundamental frequencies (cm - 1 ), solid 35C12160 at 77°K: νι(αι) 630-7, v2(a{) Λ 296-4, v 3 (W 670-8 Force constants (mdyne A~i): d . e fr = 2-75±0-02; frr = 0-40 ± 0 0 1 / 0 / r 2 = o-46 ± 0 0 1 ; frelr = 0-15 ± 0 0 1 1 Ultraviolet-visible spectrum Measured for vapour. Continuous absorption with three maxima Mass spectrum* Measured at — 78°C Positive ion 0 2 35C1 37Q 35C1Q 37C10 35CP5C1 Relative intensity 0 0 7 0-7 0-7 1000 33-4 5-4 Positive ion 35C137C1 37C137C1 35C1035C1 35C1037C1 37C1037C1 Relative intensity 3-6 0-5 43-5 28-3 4-7 Photoelectron spectrum^ Verticali.p. (eV): 11-02, 12-37, 12-65, 12-79, 15-90,16-65, 17-68,20-64 Dissociation enthalpies (kcal mol" 1 ) ClOCl -► ClO + Cl, Δ/f °298 = 34-2 (th/d); 32-3 ± 2 (mass spec.)« ClOCl -+ C l 2 + 0 , Δ#°298 = 40-4 (th/d) Dipole moment1 Measured for CC14 solution: 0-78±0-08D a

G. E. Herberich, R. H. Jackson and D . J. Millen, / . Chem. Soc. (A) (1966) 336. R. H. Jackson and D . J. Millen, Proc. Chem. Soc. (1959) 10. B. Beagley, A. H. Clark and T. G. Hewitt, / . Chem. Soc. (A) (1968) 658. d M.'M. Rochkind and G. C. Pimentel, / . Chem. Phys. 42 (1965) 1361; F. K. Chi and L. Andrews, / . Phys. Chem. 77 (1973) 3062; D. J. Gardiner, / . Mol. Spectroscopy 38 (1971) 476. e B. Beagley, A. H. Clark and D. W. J. Cruickshank, Chem. Comm. (1966) 458. 1 C. F . Goodeve and J. I. Wallace, Trans. Faraday Soc. 26 (1930) 254; W. Finkelnburg, H. J. Schumacher and G. Steiger, Z.phys. Chem. B15 (1931) 127. «I. P. Fisher, Trans. Faraday Soc. 64 (1968) 1852. h A. B. Cornford, D. C. Frost, F . G. Herring and C. A. McDowell, / . Chem. Phys. 55 (1971) 2820. 1 D . Sundhoff and H. J. Schumacher, Z . phys. Chem. B28 (1935) 17. b

c

C1N03 i

ClOE low tempNv F2

[cio][cio2][s3o10]

o°c N O S03/CFC13 AsR

AgF2

■[C102][AsF6]

TiBr, HOC1 TiOBr„ SCHEME 1. Reactions of dichlorine monoxide.

1366

CHLORINE, BROMINE, IODINE AND ASTATINE: A. J. DOWNS AND C. J. ADAMS

Chlorination of organic compounds by CI2O, C120 + 2RH -> 2RC1+H 2 0

follows a mixed chain reaction with OC1 and Cl as chain-propagating species; its selectivity is not as great as that of other chlorinating agents579. Dichlorine monoxide has found several minor synthetic applications, particularly where it can act as a substitute for the more dangerous CIO2. Its reaction with N2O5 is a very convenient route to C1N03 58 °, which is also formed from C120 and other nitrogen oxides: in each case the exact stoichiometry of the reaction depends on the phase in which it is conducted575, but the step leading to chlorine nitrate is C10 + N 0 2 - > C 1 N 0 3

The low-temperature fluorination of CI2O gives CIOF3, while AgF2 at 70°C readily converts it to C102F, probably via C102 581. The reactions whereby metal halides are transformed into oxyhalides probably involve metal hypochlorite intermediates; boron halides similarly afford anhydrous boric oxide575. Both SO3 and AsF 5 oxidize C120 in the formation of ionic complexes582. The contrasting polarizations of the halogen-oxygen bonds in CI2O and F 2 0 are marked in their reactions with 0 = CF 2 , which are catalysed by CsF: CF 3 0~ attacks the positively charged chlorine atom forming CF3OCI, whereas with F2O, CF3OOF is the initial product583. Dichlorine Trioxide Early reports of the preparation of C1203 by the partial reduction of chloric acid with reagents such as AS2O3 or undecylenic acid are unfounded; the products were mixtures of chlorine and chlorine dioxide576. However, the concept of Ο 2 0 3 as a transient intermediate accounts satisfactorily (i) for the induction time found in the photolytic decomposition of C102 in a static system, and (ii) for some aspects of the bleaching action of chlorous acid. Careful photolysis of chlorine dioxide at — 78°C affords, in addition to chlorine and oxygen, two products584: orange chlorine trioxide, and a dark brown solid identified as CI2O3. The latter compound, though stable at — 78°C, explodes on vaporizing. It is assigned the structure O-Cl-Cl

, containing a very weak chlorine-chlorine bond, and

its heat of formation is estimated to be +45 kcal mol - 1 . Chlorine Dioxide Precautions Discussion of the chemistry of chlorine dioxide (which was once called "chlorine peroxide") is prefaced by a warning of the risks involved in manipulating this capricious compound. The deep red liquid is explosive above — 40°C, and detonations can occur at temperatures as low as — 100°C; concentrations of the yellow-green vapour of > 10% in air 579 D . D . Tanner and N . Nychka, / . Amer. Chem. Soc. 89 (1967) 121. 580 M . Schmeisser, Inorganic Syntheses, Vol. 9 (ed. S. Y . Tyree, jun.), p. 127, McGraw-Hill (1967). 581 Y . Macheteau and J. Gillardeau, Bull. Soc. chim. France (1967) 4075. 582 c . J. Schack and D . Pilipovich, Inorg. Chem. 9 (1970) 387. 583 D . E . Gould, L. R. Anderson, D . E . Y o u n g and W. B . F o x , / . Amer. Chem. Soc. 91 (1969) 1310. 584 E . T. McHale and G. v o n Elbe, / . Amer. Chem. Soc. 8 9 (1967) 2795; / . Phys. Chem. 7 2 (1968) 1849.

THE OXIDES OF THE HALOGENS

1367

may explode under the action of heat, light or shock; spontaneous combustion can occur with organic or other readily oxidizable materials. Dilution with polyatomic gases stabilizes C102 to some extent. The compound is normally generated ready for immediate use, although small amounts may be stored in the liquid phase at temperatures below — 50°C or in CC14 solution in the dark: the storage of large quantities is strongly discouraged. Preparation Despite subsequent improvements in experimental conditions, the reaction between concentrated sulphuric acid and a solid chlorate, whereby chlorine dioxide was first dis­ covered, holds great risks of explosion; the essential process is the disproportionation of chloric acid. 3HC10 3 - * 2 C 1 0 2 + H C I O 4 + H 2 0

Most of the many alternative syntheses since devised involve either reduction of chlorates or oxidation of chlorites; only those methods affording convenient access to the compound are mentioned here. The oxidation of sodium chlorite with chlorine 2NaC102+Cl2 -> 2C102+2NaCl

is effected industrially by mixing concentrated aqueous solutions of the reagents574 while on a small scale chlorine gas, diluted with dry air, is passed through columns packed with the solid chlorite585. Reduction of KCIO3 in the laboratory is conveniently accomplished with moist oxalic acid, since the simultaneously evolved carbon oxides act as diluents for C102 575. 2 C 1 0 3 - + C2O4 2 - + 4 H + - > 2 C 1 0 2 + 2 C 0 2 + 2 H 2 0

Industrial processes for the reduction of chlorate depend on two major reagents: chloride ion and sulphur dioxide574. Using Cl ~, the ultimate reaction Cl - + 5C10 3 " + 6H + - > 6 C I O 2 + 3 H 2 0

is not feasible thermodynamically; under optimum conditions the reaction approximates to 2C1" + 2C10 3 " + 4 H + - > C l 2 + 2 C 1 0 2 + 2 H 2 0

and the product is contaminated with chlorine. 90% yields of CIO2 may be obtained, how­ ever, if sulphur dioxide is used to reduce solutions of sodium chlorate in 6-9 M sulphuric acid. The reaction of silver chlorate with chlorine at 90°C provided samples of pure CIO2 for the measurement of physical properties586. 2AgC103+Cl2 -> 2 C 1 0 2 + 0 2 + 2 A g C l

Structural chemistry The C102 molecule, with 19 electrons in its valence shell, has C2v symmetry in its 2 Βχ ground state. The bond length of 1-471 A is some 0-22 A shorter than the calculated single-bond length, while the O-Cl-O angle is only slightly smaller (ca. 2°) than the bond angle of sulphur dioxide (119-30)587. That the chlorine 3d orbitals are greatly involved in 585 R . 1. Derby and W. S. Hutchinson, Inorganic Syntheses, Vol. 4 (ed. J. C. Bailar, jun.), p. 152, McGrawHill (1953). 586 F . E. King and J. R. Partington, / . Chem. Soc. (1926) 925. 587 Y . Morino, Y. Kikuchi, S. Saito and E. Hirota,/. Mol. Spectroscopy, 13 (1964) 95.

1368

CHLORINE, BROMINE, IODINE AND ASTATINE: A. J. DOWNS AND C. J. ADAMS TABLE 47. SPECTROSCOPIC AND RELATED INVESTIGATIONS OF CHLORINE DIOXIDE

Microwave spectrum*»* Measured for 35Cli6() 2 , 37 Cli60 2 , 35C1160180 r,(Cl-0) = 1-473 ± 0 0 1 Ä; V c i - o = 117-6±l° a h Quadrupole coupling constants : e2Qqaa — —51-90 MHz ib cl e2Qqbb = +2-28 MHz ^ 0 /[—a e2Qqcc= +49-62 MHz 0«^

Clv symmetry

Electron diffraction0 C2u symmetry 0o-ci-o = 117-7± 1-7° r,(l)(Cl-0) = 1-475±0-003 A; Infrared spectrum Measured for vapour d and matrix-isolated moleculeh -1 Fundamental frequencies (cm ), vapour: 35C102 *ι(αι) 945-2 v2(al) 447-3 v 3 (W 1110-8 37C102 940-4 444-6 1098-1 Force constants (mdyne Ä~i): e fr = 7-23; frr = - 0 - 0 2 / e /r2 = 0-63; fjr = 0-25 Rotational analysis ofv&biY Raman spectrum Measured for solution in waters Ultraviolet-visible spectrum (absorption) Solution in H 2 0 :» Amax = 345 m^, e ~ 103 Vapour: band system 2500-5000 Ä, 2 v4 2 «- 2 2?i; vibration-rotation analysis yields excited state para­ m e t e r s ^ r(Cl-O) = 1-619±0-016Ä, V c i - o = 107-0 + 0-28°, fr = 3-69, frr = 0 0 8 , /e/r* = 0-29, / r ö /r = 0-10 mdyne Ä - i . In the vacuum-ultraviolet region three band systems are observed: at 1829 A, 1628 A and 1568 A.1 Mass spectrum* Measured at—78°C Positive ion 0 2 35Q 37Q 35QO 37C10 35C102 37C102 Relative intensity 2 1 1-9 0-6 31-3 101 1000 32-0 Ionization potential, by electron impact 10-7 ±0-1 eV l 1 Bond dissociation enthalpy (kcal mol" ) OCIO -> OCl + O, AH°29S 58-7 (th/d); 66-5 (spectroscopic) OCIO -> Cl + 0 2 , Δ # ° 2 9 8 4 0 (th/d); 4-6 (spectroscopic) Esr spectrum Measured for solutions and for the matrix-isolated molecule. In H 2 S 0 4 at 300°Kra g&v = 20093, Also» = +16-5 G In H 2 S 0 4 at 77°Km g&v = 20102, ^ l s o u = +18-0 G In KCIO3 at 295°Km g&v = 20101, A s o u = + 1 8 0 G . Other studies: n-r. Paramagnetic susceptibility3 1310xl0~6 Cgs units Measured for CCI4 solutions; 1-69±0-092) Dipole momentt a R. F. Curl, jun., J. L. Kinsey, J. G. Baker, J. C. Baird, G. R. Bird, R. F. Heidelberg, T. M. Sugden, D. R. Jenkins and C. N. Kenney, Phys. Rev. 121 (1961) 1119; R. F. Curl, jun., R. F. Heidelberg and J. L. Kinsey, ibid. 125 (1962) 1993. b R. F. Curl, jun., / . Chem. Phys. 37 (1962) 1779. c A. H. Clark and B. Beagley, / . Chem. Soc. (A) (1970) 46. d A. W. Richardson, R. W. Redding and J. C. D. Brand, / . Mol. Spectroscopy, 29 (1969) 93. e J. C. D. Brand, R. W. Redding and A. W. Richardson, / . Mol. Spectroscopy, 34 (1970) 399. f A. W. Richardson, / . Mol. Spectroscopy, 35 (1970) 43. « T. G. Kujumzelis, Physik. Z. 39 (1938) 665. h A. Arkell and I. Schwager, / . Amer. Chem. Soc. 89 (1967) 5999. 1 N. Konopik, J. Derkosch and E. Berger, Monatsh. Chem. 84 (1953) 214. 3 C. M. Humphries, A. D. Walsh and P. A. Warsop, Discuss. Faraday Soc. 35 (1963) 137,230. k I. P. Fisher, Trans. Faraday Soc. 63 (1967) 684. 1 V. I. Vedeneyev, R. V. Gurvich, V. N. Kondrat'yev, V. A. Medvedev and Ye. L. Frankevich, Bond Energies, Ionization Potentials and Electron Affinities, p. 78, Edward Arnold, London (1966). m R. S. Eachus, P. R. Edwards, S. Subramanian and M. C. R. Symons, / . Chem. Soc. (A) (1968) 1704. n P. W. Atkins, J. A. Brivati, N. Keen, M. C. R. Symons and P. A. Trevalion, / . Chem. Soc. (1962) 4785. 0 J. R. Byberg, S. J. K. Jensen and L. T. Muus, / . Chem. Phys. 46 (1967) 131. p J. C. Fayet, C. Pariset and B. Thieblemont, Compt. rend. 268B (1969) 1317. q J. E. Bennett and D. J. E. Ingram, Phil. Mag. 1 (1956) 109; J. E. Bennett, D. J. E. Ingram and D . Schonland, Proc. Phys. Soc. 69A (1956) 556. r T. Cole, Proc. Nat. Acad. Sei. 46 (1960) 506. 8 N. W. Taylor, / . Amer. Chem. Soc. 48 (1926) 854. 1 D. Sundhoff and H. J. Schumacher, Z. phys. Chem. B28 (1935) 17. u Hyperfine data for 35C1.

1369

THE OXIDES OF THE HALOGENS

the bonding is supported by molecular orbital calculations565»588; the short bond length and high valence force constant similarly testify to the strength of the chlorine-oxygen bond. The spectroscopic properties of the molecule (Table 47) are consistent with the delocalization of the unpaired electron over all three atoms in a Ab\ antibonding orbital comprising the chlorine 3p and oxygen 2p orbitals perpendicular to the plane of the mole­ cule; esr measurements indicate a spin density of 64% on chlorine589. Chlorine dioxide presents an interesting example of an odd-electron molecule which is stable towards dimerization; the physical and spectroscopic properties of the compound show no evidence of aggregation in the vapour, liquid or solid phases or in solution. This reluctance to dimerize probably stems from the delocalization of the unpaired electron. The energy expended in reorganizing the CIO2 unit to permit localization of the electron would be insufficiently offset by chlorine-chlorine overlap, which is notably weak in C1C10 and C1203. Curiously, though, the isoelectronic thionite ion (S0 2 ~) exists as S 2 0 4 2 - in crystalline Na 2 S 2 0 4 (but with a long S-S bond, 2-389 Ä) 590 , and chlorine trioxide is asso­ ciated in the liquid state despite the fact that the unpaired electron is here more delocalized than in C102. There is also some evidence that the Br0 2 radical dimerizes in aqueous solution (p. 1377). Reactions Photochemical or thermal decomposition of C102 commences with the reaction C10 2 -> C l O + O

Δ#°298 = 66-5 kcal mol"i

If the molecule is photolysed with uv light when isolated at low temperature in a rigid inert matrix589'591 (whether a solidified noble gas, a solution frozen at 77°K, or a host perchlorate lattice), the major product is the radical ClOO formed in a cage back-reaction; some ClO may also be able to flee the cage. After a short induction period, photolysis of the dry gas at room temperature proceeds with the formation of chlorine, oxygen and some chlorine trioxide which adheres to the walls of the vessel and is broken down on prolonged illumination573. CIO2 + O

->C103

cio+cio->ei2+o2 2C10 3 - * C1 2 0 6

(a)

2C103->Cl2 + 302

(b)

Since step (b) has the higher heat of activation, more chlorine trioxide is formed at lower temperatures. Photolysis of solid C102 at -78°C produces some C1203 as well as C1206 (p. 1366)584; the formation and decomposition of the sesquioxide are held to be responsible for the initial period of induction in the decomposition of C102.

cio2+cio^ci2o3 Dark green aqueous solutions containing up to 8 g C102 per litre (heat of solution = 6-6 kcal mol _1 ) decompose only very slowly in the dark; the crystalline hydrate of approximate 588 j . L . G o l e a n d E . F . H a y e s , Internat. J. Quantum Chem., S y m p o s i u m N o . 3 , p . 519 (1969-70). 589 p . w . A t k i n s , J. A . Brivati, N . K e e n , M . C . R . S y m o n s a n d P . A . Trevalion, / . Chem. Soc. (1962) 4785. 590 J. D . D u n i t z , Acta Cryst. 9 (1956) 579. 591 A. Arkell and I. Schwager, / . Amer. Chem. Soc. 89 (1967) 5999.

1370

CHLORINE, BROMINE, IODINE AND ASTATINE: A. J. DOWNS AND C. J. ADAMS

composition C102,8H20, formed when the gas is passed into cold water, is a clathrate of the 8-gas, 46-water moles type with incomplete occupation of the cells by chlorine dioxide592. Illumination of neutral aqueous solutions inaugurates rapid decomposition to a mixture of chloric and hydrochloric acids, C10 2 ->cio+o C10+H20 -> H2C102 H 2 C10 2 +C10 -►HC1+HC103

while photodecomposition of moist CIO2 gas produces a mist of droplets containing HOC1, HC102, HCIO3 and HCIO4 via hydrolysis of the intermediates C1203 and Cl20(5. In alkaline solution hydrolysis rapidly (and possibly explosively) generates a mixture of chlorite and chlorate anions, and for this reason C102 is regarded as the mixed anhydride of chlorous and chloric acids. C102,8H20(±H20) M(CKX) (n = 2,M = Mg,Cd,Zn,Ni;

C102,MF5 (M = As, Sb)'

cio3~ + α ο , -

Na[NH,ClO]

CINO,

(cio)(cio2)s3oip <™>2>2S3°,0 SCHEME 2. Reactions of chlorine dioxide.

Some reactions of chlorine dioxide are summarized in Scheme 2. The compound is a strong oxidizing agent, combining readily with organic matter and numerous inorganic materials, including phosphorus, sulphur, phosphorus halides and potassium borohydride. The reaction of CIO2 with iodide ion involves CIO2I" as intermediate593; the stoichiometry of the reaction depends on the pH of the solution. Acid solution: Neutral solution:

2ClO2+10I-+8H+ 2C102+2I-

► 2C1-+4H20+5I2 2C102-+I2

592 M. Bigorgne, Compt. rend. 236 (1953) 1966. 593 H . Fukutomi and G. Gordon, / . Amer. Chem. Soc. 89 (1967) 1362.

THE OXIDES OF THE HALOGENS

1371

The latter reaction suggests that, in the presence of a suitable reductant, the disproportionative hydrolysis of CIO2 in alkaline solution may be diverted to the production of chlorites; the most convenient reagents for this purpose are alkaline peroxides. 2CIO2+O2 2 - - * 2 C 1 0 2 + 0 2

The oxidation of powdered metals to chlorites by cold aqueous solutions of CIO2 probably involves direct electron transfer; for the redox potential of the process C10 2 (aq)+e-(aq) ^ C10 2 -(aq) 572

electrochemical measurements give a value of +0-936 V at 25°C (cf. Fig. 14 and ref. 289). Fluorination of C102 with fluorine at -50°C or AgF 2 at 20°C readily affords chloryl fluoride575. The reaction of gaseous CIO2 with fluorine is homogeneous and bimolecular at low partial pressures and low temperatures594, the rate-determining step being C102+F2->C102F+F

At room temperature only spontaneous decomposition of C102 is observed when fluorina­ tion is attempted. Little is known about the 1:1 adducts with arsenic or antimony pentafluoride, which are stable at temperatures up to 80°C595; the nature of sodium amidochlorate Na[NH2C102] is equally obscure596. Analysis514 The most convenient method of determining chlorine dioxide is by liberating iodine from neutral or acidic iodide solution, followed by thiosulphate titration. The presence of halogen oxyanions causes complications, since these species also oxidize I - to I2; chlorine dioxide may be blown out of solution in a current of nitrogen, and absorbed in neutral KI solution. C102 can be estimated in the presence of chlorine by treatment with acid KI and back-titration to determine the acid consumed. C/^574,597

Chlorine dioxide, manufactured by the reduction of chlorates, is used in making sodium chlorite (subsequently employed for bleaching), and is itself used for the bleaching of paper pulp on the scale of 105 tons per annum; the bleaching and oxidizing powers of the com­ pound have also been applied to oils, fats, waxes, flour and textiles, to treating leather prior to tanning (by decomposing the keratin), to improving the viscosity of rubber-based varnishes and glues, and to sterilizing foodstuffs such as cottage cheese. When used in the purification of water, chlorine dioxide is normally generated by the action of chlorine on sodium chlorite. CIO2 enjoys several advantages over chlorine in its ability to destroy ill-tasting phenols and to convert iron and manganese rapidly to insoluble forms. Chlorine Perchlorate (see also p. 1473) This unusual chlorine oxide598 is prepared by the reaction MCIO4+ C10S0 2 F

C

> MSO3F+ CIOCIO3

(M = Cs or N 0 2 )

594 p . j . A y m o n i n o , J. E . Sicre and H . J. Schumacher, / . Chem. Phys. 2 2 (1954) 756. 595 M . Schmeisser and W . Fink, Angew. Chem. 6 9 (1957) 7 8 0 . 596 G . Beck, Z. anorg. Chem. 2 3 3 (1937) 155. 597 w . Masschelein, Chim. Ind., Genie Chim. 9 7 (1967) 4 9 , 3 4 6 . 598 c . J. Schack and D . Pilipovich, Inorg. Chem. 9 (1970) 1387.

1372

CHLORINE, BROMINE, IODINE AND ASTATINE: A. J. DOWNS AND C. J. ADAMS

It is thermodynamically and kinetically less stable than chlorine dioxide, decomposing at ambient temperatures to CI2, O2 and C^Og. The formally positive chlorine atom is quanti­ tatively replaced by bromine in forming the mixed oxide BrOC103 5 " : Br2 + 2CIOCIO3 — °-> Cl 2 + 2BrOC10 3

Dichlorine Hexoxide Preparation The best method of making C1206 is probably the ozonolysis of chlorine dioxide600. 2C10 2 + 2 0 3 - ^ C l 2 0 6 + 2 0 2

The compound is also produced in the photochemical combination of chlorine and ozone, in the photolysis of chlorine dioxide, and (with 80% yield) in the decomposition of chlorine perchlorate598. A neat and original synthesis involves the thermal degradation of the diperchlorate of xenon601. 20°

Xe(C10 4 ) 2

^ X e + 0 2 + Cl 2 0 6

Structure In the vapour phase the compound probably exists as the pyramidal CIO3 molecule602; this paramagnetic species has also been studied by esr and ultraviolet-visible absorption spectroscopy, both in frozen solutions and trapped in crystalline lattices (Table 48). Along the isoelectronic series PO32 -, S0 3 ~, CIO3, the esr parameters indicate increasing delocalization of the odd electron onto the oxygen atoms; there is an accompanying increase in bond angle: P 0 3 2 - , 110°; S0 3 ~, 111°; C103,112° 6°3. Photolysis of the isolated C103 molecule readily produces ClO and oxygen604. hv

cio3->cio+o2

It should be noted that the ultraviolet-visible spectrum attributed to gaseous CIO3 605a (Table 48) bears little resemblance to that for CIO3 produced by the y-radiolysis of 12-5 M HCIO4 frozen at - 196°C604; the vapour-phase spectrum, however, is distinctly like that of ClO (p. 1381). The molecular weight of the compound, measured cryoscopically in carbon tetrachloride, is consistent with a dimeric formulation Ο2θ 6 , as are the properties of solutions in oleum and H3PO4. y-Radiolysis of solutions of CI2O6 frozen at - 196°C produces C10 3 ; photol­ ysis of such solutions gives ClO as the principal decomposition fragment604. In the liquid and solid phases the oxide is said to exist exclusively as the dimer Ο2θ 6 , although the high melting point and boiling point could indicate a polymeric formulation. Early measurements of the paramagnetic susceptibility of the liquid605a were interpreted in terms of the equilibrium C1 2 0 6 ^ 2C10 3 599 C. J. Schack, K. O. Christe, D. Pilipovich and R. D. Wilson, Inorg. Chem.10 (1971) 1078; K. O. Christe and C. J. Schack, ibid. 13 (1974) 1452. 600 H. J. Schumacher and G. Stieger, Z. anorg. Chem. 184 (1929) 272. 601 N . Bartlett, M. Wechsberg, F. O. Sladky, P. A. Bulliner, G. R. Jones and R. D. Burbank, Chem. Comm. (1969) 703. 602 c . F. Goodeve and F. A. Todd, Nature, 132 (1933) 514. 603 A. Begum, S. Subramanian and M. C. R. Symons, / . Chem. Soc. (A) (1970) 918. 604 v . N. Belevskii and L. T. Bugaenko, Russ. J. Inorg. Chem. 12 (1967) 1203. 605 ( a ) J. Farquharson, C. F. Goodeve and F. D. Richardson, Trans. Faraday Soc. 32 (1936) 790; (b) A. Pavia, J. L. Pascal and A. Potier, Compt. rend. 272C (1971) 1495.

THE OXIDES OF THE HALOGENS

1373

TABLE 48. SPECTROSCOPIC AND RELATED INVESTIGATIONS OF CHLORINE TRIOXIDE/DICHLORINE HEXOXIDE

Ultraviolet-visible spectrum CIOz Vapour:» A m a x ~ 2 7 8 0 Ä , e ~ 1 ·2 χ 10* (see text) In 12-5 M HC10 4 at - 196°C: b Amax — 4300 A, € ~ 5 x 103 ClzOe Liquid;» CCI4 solution; c oleum solution.1* Magnetic susceptibility* Measured using Gouy method as a function of temperature between — 40°C and + 10°C. For the dissociation C1 2 0 6 ^ 2C10 3 , log K = -0-974-1730/2-3RT(see text) Esr spectrum Measured for CIO3 in frozen solutions or isolated in crystalline lattices; the radical is produced by y-radiolysis or X-ray irradiation of an oxychlorine substrate. In 12-5 M HCIO4 at 77°K: d g&v = 2009, Alao e = 128-6 G In KCIO4 at 300°K: f gAV = 2 0 1 1 , Alao e = 122 G In NH4CIO4 at 300°K:e g&v = 2 0 0 8 , Aiso e = 128 G Also for CIO3 in LiC10 4 (195°K),f NaC10 4 (195°Κ),< Mg(C10 4 ) 2 (195°K),f aq. HCIO3 (77°K), d C1 2 0 6 (77°K),b oleum (77°K),b FC10 4 (77°K),b N0C10 4 (77°K),b NaC10 3 (25-298°K),h-i NaC10 3 and KC1Q3 (26°K)J a

C. F. Goodeve and F. D. Richardson, Trans. Faraday Soc. 32 (1936) 790. V. N. Belevskii and L. T. Bugaenko, Russ. J. Inorg. Chem. 12 (1967) 1203. « M. H. Kalina and J. W. T. Spinks, Canad. J. Res. B16 (1938) 381. * V. N . Belevskii and L. T. Bugaenko, Zhur. Fiz. Khim. 41 (1967) 144. e Hyperfine data for 35C1. f P. W. Atkins, J. A. Brivati, N . Keen, M. C. R. Symons and P. A. Trevalion, / . Chem. Soc. (1962) 4785. * Τ. Cole, / . Chem. Phys. 35 (1961) 1169. h F. T. Gamble,/. Chem.Phys. 42(1965) 3542; J. C. Fayet and B. Thieblemont, Compt. rend. 261 (1965) 1501. 1 O. Vinther, / . Chem. Phys. 57 (1972) 183. j J. R. Byberg, Chem. Phys. Letters, 23 (1973) 414. b

but recent esr investigations of the condensed phases could not detect the presence of CIO3604; the sole paramagnetic species was the impurity C102 in 0Ό1 M concentration. Little credence can therefore be given to the much-quoted heat of dissociation, viz. 1 -73 kcal mol -1 . No structural data are available for the dimer, but there are two outstanding possibilities, viz.

/°

Cl—O

\ (4)

\ A /° / V \ Cl

Cl

(5)

Structure (4) is attractive in that little rearrangement of the C103 pyramids is needed to pair the odd electrons, although the bridged dimer (5) comes closer to the ionic formulation C102 + C104", which is judged to be compatible with the vibrational spectrum attributed to C1206 at 87°K605*. Reactions The dark red liquid is allegedly the least explosive of the chlorine oxides. In many of its reactions it appears to behave as chloryl perchlorate [C102]+[C104] - 575 ; similar behaviour is exhibited, for example, by diborane in some of its reactions. Nitrogen oxides and their derivatives displace C102 with formation of nitrosyl or nitryl perchlorates though the reactions are not reversible. With anhydrous HF, an equilibrium is set up and C102F and HC104 are formed; C1206 may also be used to synthesize metal perchlorates. With AsF5 an unusual CIO3 adduct of unknown structure is formed. Hydrolysis gives a mixture of chloric and perchloric acids. C 1 2 0 6 + H 2 0 -+ HCIO3+HCIO4

1374

CHLORINE, BROMINE, IODINE AND ASTATINE: A. J. DOWNS AND C. J. ADAMS (NqpClO, + C102

(NOx)C104 NQX (x =1,2)

N

+ cio 2 + 1/2 Cl2

Hp

sNOxCl

(*- ■1.2) CI

2°6

-^Cr03

Cr02Cl2 or Cr02(C104)2

^sr

^^HC104 + C102F

-iopcVF5

soci2

C103,AsF5

(cio2); S3°,0 SCHEME 3. Reactions of dichlorine hexoxide.

Dichlorine Heptoxide Preparation The oily liquid is best obtained by the careful dehydration of perchloric acid with phosphoric acid at — 10°C, followed by equally careful distillation at - 3 5 ° C a n d l m m pressure606. It is also formed when chlorine and ozone react in blue light. Structure The gaseous molecule has a O3CI-O-CIO3 structure of C2 symmetry, the CIO3 groups beiiig 15° from the symmetrically staggered C2V configuration. The dimensions of the bridging system (Table 49) are consistent with "normal" single bonds; repulsion between the CIO3 groups accounts for the opening of the Cl-O-Cl angle from the expected tetrahedral value. Whether the shortness of the terminal Cl-O bonds evinces d„-p„ bonding was dis­ cussed in Section 4B1. Reactions CI2O7 is relatively stable for a chlorine oxide; though exploding when heated or sub­ jected to shock, it does not ignite organic material at room temperature. Its synthesis shows it to be the true anhydride of perchloric acid, and it regenerates CIO4 ~ on dissolution in water or alkali. Thermal decomposition of dichlorine heptoxide607 to chlorine and oxygen in both the liquid and vapour phase is initiated by rupture of a bridging Cl-O bond: C1 2 0 7 -*C103+C104

The activation energy for this process is 32-9 kcal mol - 1 for the vapour and 32-1 kcal mol -1 for the liquid. 606 c . F. Goodeve and J. Powney, / . Chem. Soc. (1932) 2078. 6071. p. Fisher, Trans. Faraday Soc. 64 (1968) 1852.

1375

THE OXIDES OF THE HALOGENS TABLE 49. SPECTROSCOPIC AND RELATED INVESTIGATIONS OF DICHLORINE HEPTOXIDE

Electron diffraction* Molecule has O3CIOCIO3 structure of C2 symmetry r,(l)[Cl-Oterm] = 1-405 ± 0 0 0 2 A; r f (l)[Cl-O br ] = 1-709±0004Ä 0ci-obr-ci = 118·6±0·7°; 0ot-ci-ot = Π5·2±0·2° CIO3 groups are oriented 15° from Civ (staggered) configuration. Cl-Obr bonds are inclined at 4-7° to threefold axis of CIO3 groups. infrared spectrum* Measured for vapour and solid. Raman spectrum0 Measured for liquid. Force constants* /r(Cl-Oterm) = 9-32 mdyne A"i / r (Cl-O br ) = 3-2mdyneA-i Ultraviolet absorption spectrum* Measured for vapour. Dipole moment1 Measured for CC14 solution: 0-72±0-02D. Mass spectrum* 35C1 37d 35dO Positive ion O2 Relative intensity 0-27 901 108 0-34 35 35C137C1 37d 2 Positive ion Cl2 35d03 Relative intensity 0-99 0-54 011 1000 35 Positive ion Cl2<>7 35C137C107 37C1207 Relative intensity 0-67 0-45 009 Dissociation enthalpy O3CI-OCIO3 -> O3CI+CIO4; ΔΗ = 30 ± 4 kcal mol" 1

37C10 2-88 37C103 33-4

35C102 31-2 35C104 0-67

37C102 10-4 37C104 0-23

» B. Beagley, Trans. Faraday Soc. 61 (1965) 1821. * R. Savoie and P. A. Giguöre, Canad. J. Chem. 40 (1962) 991. c J. D. Witt and R. M. Hammaker, Chem. Comm. (1970) 667. d Based on incorrect assignment: E. A. Robinson, Canad. J. Chem. 41 (1963) 3021; R. J. Gillespie and . A. Robinson, ibid. 42 (1964) 2496. e C. F. Goodeve and B. A. M. Windsor, Trans. Faraday Soc. 32 (1936) 1518. f R. Fonteyne, Natuurw. Tijdschr. 20 (1938) 112, 275. * I. P. Fisher, Trans. Faraday Soc. 64 (1968) 1852.

Chlorine Tetroxide Gomberg claimed that the reaction of iodine with silver perchlorate in anhydrous ether produced chlorine tetroxide608. I 2 +2AgC10 4 -> 2AgH-(C10 4 ) 2

Subsequent investigations of this reaction have shown the products to be iodine perchlorates (p. 1473)609. The CIO4 molecule is a probable intermediate in the thermal decomposition of CI2O7 607 . A species of this stoichiometry is also formed in the y-radiolysis of potassium chlorate at 77°K, but probably has the structure

\

Cl-O-O^io.

The electron affinity of the hypothetical molecule is calculated to be 134 kcal mol _ 1 (from the lattice energies of perchlorates)611. 608 M. Gomberg, / . Amer. Chem. Soc. 45 (1923) 398. 609 N. W. Alcock and T. C. Waddington, / . Chem. Soc. (1962) 2510. 610 R. s. Eachus, P. R. Edwards, S. Subramanian and M. C. R. Symons, / . Chem. Soc. (A) (1968) 1704. 611 V. I. Vedeneyev, L. V. Gurvich, V. N. Kondrat'yev, V. A. Medvedev and Ye. L. Frankevich, Bond Energies, lonization Potentials and Electron Affinities, Edward Arnold, London (1966).

1376

CHLORINE, BROMINE, IODINE AND ASTATINE: A. J. DOWNS AND C. J. ADAMS

BROMINE OXIDES575.612.613 Dibromine Monoxide In accord with the synthesis of C120, Br 2 0 is formed in the reaction of mercuric oxide with bromine vapour or with cold solutions of bromine in carbon tetrachloride. 2Br2 + 2HgO -> HgBr2.HgO+Br20 Yields are low, particularly for the vapour-phase reaction, and pure dibromine monoxide is best prepared by the decomposition of Br0 2 in vacuo614; pumping off other products at -60°C leaves Br 2 0. TABLE 50. PROPERTIES OF DIBROMINE MONOXIDE

Colour of solid:* brown-black Melting point:a — 17 ·5°C (with decomposition) Infrared spectrum:* Measured for solid at — 196°C Fundamental frequencies (cm -1 ) and assignments in C2„ symmetry. VI(ÖI) 504; ρ2(αι) 197; v3(6i) 587. Force constants (mdyne A"i) fT = 2·4±0·2; f„ = 0·4±0·2; / > = 0·2±0·1. 0 Ultraviolet-visible spectrum Measured for solutions in CCI4.

/e/r* = 0·4±0·1;

* M. Schmeisser and K. Brändle, Adv. Inorg. Chem. Radiochem. 5 (1963) 41. C. Campbell, J. P. M. Jones and J. J. Turner, Chem. Comm. (1968) 888. W. Benschede and H. J. Schumacher, Z. anorg. Chem. 226 (1936) 370.

b c

The compound is unstable above — 40°C with respect to decomposition into bromine and oxygen. The infrared spectrum of the brown-black solid at - 196°C is consistent with a bent C2v structure for the BrOBr molecule; the estimated bond angle is 11Γ 6 1 4 . Few reactions of the compound have been studied.

OBr~

ϊ2ο5

SCHEME 4. Reactions of dibromine monoxide. Bromine Dioxide 575 »613 The quantitative ozonolysis of bromine in a fluorocarbon solvent at low temperatures produces bromine dioxide as a light yellow crystalline solid unstable much above — 40°C. Br 2 +40 3 612

CF3C1

e -78 C

> 2Br0 2 +40 2

A. G. Sharpe, Mellows Comprehensive Treatise on Inorganic and Theoretical Chemistry; Supplement II, Part I, pp. 747-749, Longmans, London (1956). 613 P. J. M. Radford and M. Schmeisser, Bromine and its Compounds (ed. Z. E. Jolles), p. 147, Benn, London (1966). «* C. Campbell, J. P. M. Jones and J. J. Turner, Chem. Comm. (1968) 888.

THE OXIDES OF THE HALOGENS

Q

iodine

O Oxygen

FIG. 31. Crystal structure of I2O5.

TABLE 52. SOME PROPERTIES OF IODINE PENTOXIDE

Ai//°(s) at 298°K Crystal structure*

Intermodular interactions are depicted in Fig. 31.

-37-78 kcal mol-i a Intramolecular distances (A) I1-O1 1-78(3) I2-O3 I1-O2 1-77(3) I2-O4 I1-O5 1-92(2) I2-O5

1-83(3) 1-79(3) 1-95(3)

Intramolecular angles (°) O1-I1-O2 99-5(1-3) O1-I1-O5 96-5(1-2) O2-I1-O5 101-9(1-0) O3-I2-O4 94-8(1-1) O3-I2-O5 931(11) O4-I2-O5 97-5(1-0) I1-O5-I2 139-2(1-4) Measured for solid I2O5

Infrared and Raman spectra0 Mass spectrum* 127/ Quadrupole resonance spectrum^ Measured at 77°K Quadrupole coupling constant (77°K) e^Qq = 1068 MHz η = 0-33

a Selected Values of Chemical Thermodynamic Properties, p. 37, N.B.S. Technical Note 270-3 (1968). b K. Seite and A. Kjekshus, Acta Chem. Scand. 24 (1970) 1912. c P. M. A. Sherwood and J. J. Turner, Spectrochim. Acta, 26A (1970) 1975. d M. H. Studier and J. L. Huston, / . Phys. Chem. 71 (1967) 457. e S. Kojima, K. Tsukada, S. Ogawa and A. Shimauchi, / . Chem. Phys. 23 (1955) 1963.

1380

CHLORINE, BROMINE, IODINE AND ASTATINE: A. J. DOWNS AND C. J. ADAMS

contains a network of interacting I2O5 and HIO3 molecules620; the I 2 0 5 bond lengths match those found in the pure pentoxide, but the I-O-I bridging angle is much larger in I 2 0 5 (139-2°) than in anhydriodic acid, and the relative orientations of the I 0 2 groups are also different. I 2 0 5 oxidizes many common materials such as NO, C2H4, H2S and CO: the last of these reactions proceeds quantitatively at ambient temperatures when CO is passed over the powdered oxide617. i 2 05 + 5 C O - > I 2 + 5C02

The ease with which iodine may be determined makes this a very useful method for the estimation of CO in the atmosphere or in other gaseous mixtures. Fluorination of I 2 0 5 (with F 2 , BrF3, C1F3, SF4 or C102F) affords IF 5 , which itself reacts with I 2 0 5 to give IOF 3 . With S0 3 or S 2 0 6 F 2 , salts of the iodyl cation I 0 2 + are formed621, while in concentrated acids (H2SO4, H2S2O7, H2Se04) iodine pentoxide is reduced by iodine to iodosyl derivatives (IO)2X (X = S0 4 , S 2 0 7 , Se0 4 ) 622 . Other Oxides The I 0 2 radical (half-life 50 ^sec) has been detected in flash photolysis of aqueous iodate solutions (p. 1382), but on the testimony of its infrared and Mössbauer spectra the yellow diamagnetic crystalline solid "l 2 0 4 "is considered to be iodosyl iodate [IO] + [I0 3 ] ~ 623 ; the iodosyl cation presumably forms a polymeric chain [(I-0-) + ], as in other iodosyl derivatives, and is cross-linked by iodate anions. Alkaline hydrolysis of I2O4 produces iodate and iodide. 3 I 2 0 4 + 6OH - -> 5 I 0 3 " + 1 - + 3 H 2 0

Hydrochloric acid degrades it to iodine monochloride. I2O4+8HCI-+2ICI+3CI2+4H2O

The commonly held view that the yellow hygroscopic material I 4 0 9 is in fact iodine triiodate I(I0 3 ) 3 is based solely on the apparent stoichiometry of its reactions with water and with hydrogen chloride575. The existence of other iodine oxides is uncertain. By analogy with the synthesis of CI2O7, attempts to prepare I 2 0 7 have centred on the dehydration of periodic acid624: the most convincing claim is that 65% oleum produces a highly reactive orange solid with approximately the correct composition625. Thermal dehydration of paraperiodic acid at ca. 130°C produces a solid I2O5J2O7. Diiodine trioxide is unknown, but the iodosyl cation (p. 1352) and various iodosyl salts are derivatives of this hypothetical oxide. Oxyhalogen Radicals As well as the oxides which may be obtained as pure materials, a number of molecular species have been identified, either as transients in the vapour phase or in solution, or trapped in solid matrices. 620 Y . D . F e i k e m a a n d A . Vos, Ada Cryst. 2 0 (1966) 769. 621 F . A u b k e , G . H . C a d y a n d C . H . L . K e n n a r d , Inorg. Chem. 3 (1964) 1799. 622 G . D a e h l i e a n d A . Kjekshus, Acta Chem. Scand. 18 (1964) 144. 623 j . H . Wise a n d H . H . H a n n a n , / . Inorg. Nuclear Chem. 2 3 (1961) 3 1 . 624 M . D r ä t o v s k y a n d L . Pacesovä, Russ. Chem. Rev. 37 (1968) 2 4 3 . 625 H . C. Mishra a n d M . C . R. Symons, / . Chem. Soc. (1962) 1194; H . Siebert a n d G . Wieghardt, Z . Naturforsch. 27b (1972) 1299; H . Siebert and U . Woerner, Z . anorg. Chem. 398 (1973) 193.

1377

THE OXIDES OF THE HALOGENS

The compound also results from the action of oxygen atoms on bromine, e.g. during glowdischarge of bromine-oxygen mixtures. On the evidence of its Raman spectrum6153, the solid may consist of dimeric units 0 2 Br· Br0 2 . Calorimetric study of the violent decomposition which occurs on rapid warming to 0°C, 2Br02-+Br2+202 implies that A///°[Br02(s)] is 4-12-5 ±0-7 kcal mol - 1 ; the compound is less endothermic than C102 (ΔΗ/° = 25-1 kcal mol" 1 ). Slow warming of Br0 2 under vacuum evolves Br 2 0 and a white solid—presumably a higher oxide. Hydrolysis of bromine dioxide, which proceeds in 5 M alkali, reflects the instability of Br(III) in aqueous solution. 6 B r 0 2 + 6 0 H - -> 5Br0 3 " + B r " + 3 H 2 0

Reactions with N 2 0 4 and fluorine have also been investigated. [N0 2 ] + [Br(N0 3 ) 2 ]" ^

Br02 - ^ * Br02F

About the discrete BrC>2 radical little is known. It may be generated in aqueous solution by the flash photolysis or pulse radiolysis of Br0 3 - (pp. 1380-5); the kinetics of its hydrol­ ysis by OH~ suggest that it may exist in equilibrium with the dimer Br 2 0 4 6 1 5 b . 2Br02^Br204

tf

= 19M"1

Higher Oxides The higher oxides of bromine have not been well characterized. While the ozonolysis of bromine to Br0 2 proceeds smoothly at — 78°C in an inert solvent, execution of the reaction in the vapour phase at higher temperatures has yielded at least three products, depending on the material of the reaction vessel and the pressure of the reactants: only approximate elemental analyses have commonly been given. Any (or all) of these white solids may also be formed in the controlled decomposition of Br(>2. No structural studies of these com­ pounds have been reported, and attempts to dehydrate perbromic acid to Br 2 0 7 have been unsuccessful. TABLE 51. THE HIGHER OXIDES OF BROMINE»»1»

Composition Colour Preparation

Br 2 O s white Ozonolysis of Br2 at 0°C in pyrex: [Br2]

Hydrolysis product

HBr0 3

Br 3 O s white Ozonolysis of Br2 at 0°C in quartz:

5

a

<-£T<15 [Br2]

"H 4 Br 3 O 10 "

b

Br0 3 white Ozonolysis of Br2 at 0°C: [Br2] Glow discharge on 2% mixtures of Br2 in 0 2 HBr03+02

* M. Schmeisser and K. Brändle, Adv. Inorg. Chem. Radiochem. 5 (1963) 72. P. J. M. Radford and M. Schmeisser, Bromine and its Compounds (ed. Z. E. Jolles), p. 147, Benn, London (1966). b

615

(a) J.-L. Pascal and J. Potier, Chem. Comm. (1973) 446; (b) G. V. Buxton and F. S. Dainton, Proc. Roy. Soc. A304 (1968) 427.

1378

CHLORINE, BROMINE, IODINE AND ASTATINE: A. J. DOWNS AND C. J. ADAMS

The Br0 3 radical, produced by the y-radiolysis of KBr0 3 at 77°K, has been studied by esr spectroscopy (Table 54); the O-Br-O angle is estimated to be 114° 603. IODINE

l

OXIDES«5,576,6i6

]

Λ

2

(or IBr)

h°<

HIO

75 °C

135°C

3/2 I 2 0 5 + l / 2 I 2 + 3 / 4 0 2

—

4/5 I , 0 5 + 1/5 I 2

H1

A

IO,SO F (+o")

(io 2 ) 2 so 4

S,(),F,

-±±J—

i2a

IOF„

300°C =r glow discharge

i^ + 5/2 O,

(lO) 2 X ( X = S 0 4 , S e 0 4 ,

S20?)

SCHEME 5. The oxides of iodine.

Iodine Pentoxide617 I 2 0 5 is the most important and thermally the most stable of the iodine oxides. In the most convenient preparation, iodic acid is dehydrated at 200°C in a stream of dry air. The hygroscopic white solid absorbs water readily from the atmosphere, giving HI 3 0 8 (formu­ lated ΐ2θ 5 ,ΗΙθ3); commercial samples of "iodine pentoxide" may contain copious amounts of this hydrate618. On dissolution in water iodine pentoxide regenerates HI0 3 . The crystal structure of the pentoxide619 contains I 2 0 5 molecules (Fig. 31), consisting of two IO3 pyramids joined through one corner: the bridging I-O distances correspond to single bonds, whereas the terminal bonds are considerably shorter. That intermolecular iodine-oxygen contacts (^2-23 A; represented by broken lines in Fig. 31) contribute significantly to the bonding is also evident from the vibrational spectra of the solid. HI3Og 616 G. J. Hills, Mellows Comprehensive Treatise on Inorganic and Theoretical Chemistry, Supplement II, Part I, pp. 870-873, Longmans, London (1956). 617 A. W. Hart, M. G. Gergel and J. Clarke, Kirk-Othmer Encyclopedia of Chemical Technology, 2nd edn., Vol. 11, p. 859, Interscience, New York (1966). 618 K. Seite and A. Kjekshus, Acta Chem. Scand, 22 (1968) 3309. 619 K. Seite and A. Kjekshus, Acta Chem. Scand. 24 (1970) 1912.

THE OXIDES OF THE HALOGENS

1381

TABLE 53. OXYHALOGEN RADICALS

CICIO* C10 a » bc [OC10]d C!OO a b [C10 3 ] d

cio 4 a

BrO b c OBrOc BrOOb BrO 3 a 0

IObc OIOc IOOb I0 3 C

a

Characterized by matrix-isolation. Characterized in the vapour phase. Characterized in solution. d Stable species under normal conditions.

b c

ClO, BrO and 10 The radicals ClO, BrO and IO are produced in the vapour phase (i) by the introduction of the halogen or its compounds into an oxyhydrogen flame, (ii) by ozonolysis of the halogen, or (iii) by the action of flash photolysis or microwave discharge on halogen-oxygen mixtures. The most important reactions are: and

o+x 2 ->xo+x x+o 2 ->xo+o

ClO is an efficient carrier in chain reactions, and its role as an intermediate in the synthesis and decomposition of other oxychlorine molecules is well exemplified in preceding sections. Detailed spectroscopic studies of the vapour-phase radicals have yielded the ground state parameters listed in Table 55. The high dissociation energy and short bond lengths of ClO are evidence of a strong bond, as is the fundamental frequency of 995 cm - 1 attributed to 35C10 isolated in an argon matrix, with the derived force constant of 6-4 mdyne A - 1 626 ; Hartree-Fock SCF-MO calculations predict a stretching frequency of 975 cm - 1 5 6 6 . (However, analyses of the electronic and microwave spectra of the gaseous molecule suggest lower stretching frequencies—868 and 610 + 60 cm™1 respectively.) The reactions of ClO, BrO and IO in the vapour phase have been studied using flow systems627. In the absence of competing species, the gaseous XO radicals decompose according to the scheme(s):

xo+xo-*x+xoo XOO -> X + 0 2 X O O + X -> X 2 + 0 2

(X = Cl, Br or I) (X = Cl)

Dissociation of the unstable peroxybromine and peroxyiodine molecules is immediate; the more stable ClOO may also degrade, however, via reaction with chlorine atoms. At high pressures decomposition of ClO may proceed through the dimer (C10)2- This dimer is also found in inert gas matrices following the photolysis of CI2O isolated at 14°K; in 626 L. Andrews and J. I. Raymond, / . Chem. Phys. 55 (1971) 3087. 627 M. A. A. Clyne and J. A. Coxon, Trans. Faraday Soc. 62 (1966) 1175; M. A. A. Clyne and H. W. Cruse, ibid. 66 (1970) 2214,2227.

ca. 295

aq. NaOH

IO3

uv, kinetics

uv, kinetics

kinetics

kinetics kinetics kinetics kinetics

n

n

j k j k,l m n

g h

esr esr uv, uv, uv, uv, esr uv,

f

d e

o

o

Amax = 3800 A

Amax = 7150 A

i*av = 2-070; Λ1β0 = 614 G AmftX = 4900Ä

Amax = 4750 A

Amax = 3500 Ä; half-life ca. 2/xsec

Probable structure / >C1—Ov x

Fundamental frequencies (cm - 1 ): 1441, 407, 373. Force constants: / r ( 0 - 0 ) , 9-65 mdyne A~i; / r (0-Cl), l-29mdyneA~i.

Weak (ρ-π*)σ Cl-Cl bond Fundamental frequency: 35 C10 994-8 c m - i ; 37QO 9860 cm"**. Force constant = 6-41 m d y n e A - 1 g„= l-9909,# v v = 20098, g„= 1·9909,*. τ «1-9972; Axx = 5-7, Ayy = - 1 1 - 4 , A„ = 5-7, AlBO= - 5 - 7 G° ^max ca. 2650 Ä Broad absorption at 2400 A

a b c

Comments

Ref.

uv ir, uv, kinetics ir

esr

ir ir

Technique

a M. M. Rochkind and G. C. Pimentel, / . Chem. Phys. 46 (1967) 4481. *° L. Andrews and J. I. Raymond, / . Chem. Phys. 55 (1971) 3087. c P. W. Atkins, J. A. Brivati, N. Keen, M. C. R. Symons and P. A. Trevalion, / . Chem. Soc. (1962) 4785. d I. Norman and G. Porter, Proc. Roy. Soc. A230 (1955) 399.

Br0 3 IO1

OBrO

ca. 295

y-radiolysis of CIO4" y-radiolysis of CIO4"

77 77

KCIO4 KCIO4

aq. NaOH

photolysis of chlorine dioxide

4

Solid argon

io2

u.v. photolysis of chlorine dioxide photolysis of CI2 in presence of O2

77 ca. 300

Solid H2SO4 Vapour

pulse y-radiolysis of OBr~, Br0 2 " flash photolysis of BK>2~ pulse radiolysis of Br0 2 ~, Br0 3 ~ flash photolysis of Br0 2 ", Br0 3 ~ y-radiolysis of Br0 3 " (at 77°K) formed in secondary reactions after pulse y-radiolysis and flash photolysis of iodine oxyanions. pulse y-radiolysis and flash photolysis of io3pulse y-radiolysis and flash photolysis of IQ3-

u.v. photolysis of chlorine dioxide

77

Solid H3PO4 and solid H2SO4

ca.295 ca.295 ca.295 ca. 295 ca.130 ca. 295

photolysis of dichlorine monoxide Cl 2 0 + Li->C10+LiCl

14 4

Solid nitrogen Solid argon

Method of production

Temp. (°K)

Medium

aq. NaOH aq. NaOH aq. NaOH aq. NaOH KBr0 3 aq. NaOH

BrO

1

C104

ClOO

cioh

CICIO

Molecule

TABLE 54. SPECTROSCOPIC STUDIES OF SOME OXYHALOGEN RADICALS

• H. S. Johnston, E. D. Morris, jun. and J. Van den Bogaerde, / . Amer. Chem. Soc. 91 (1969) 7712. A. Arkell and I. Schwager, /. Amer. Chem. Soc. 89 (1967) 5999. < R. S. Eachus, P. R. Edwards, S. Subramanian and M. C. R. Symons, /. Chem. Soc. (A) (1968) 1704. h J. R. Morton, /. Chem. Phys. 45 (1966) 1800. 1 For information about vapour-phase molecule see Table 55. J G. V. Buxton, F. S. Dainton and F. Wilkinson, Chem. Comm. (1966) 320. k O. Amichai, G. Czapski and A. Treinin, Israel J. Chem. 7 (1969) 351. 1 F. Barat, L. Gilles, B. Hickel and J. Sutton, Chem. Comm. (1969) 1485. m A. Begum, S. Subramanian and M. C. R. Symons, / . Chem. Soc. (A) (1970) 918. n O. Amichai and A. Treinin, / . Phys. Chem. 74 (1970) 830. 0 35 C1 hyperfine interactions.

f

1384

CHLORINE, BROMINE, IODINE AND ASTATINE! A. J. DOWNS AND C. J. ADAMS TABLE 55. PROPERTIES OF THE VAPOUR-PHASE SPECIES ClO, BrO and IO

Thermodynamic properties at 298°K Δ7/, 0 (kcal mol "i) a AG,° (kcal mol"*)a ^(caldeg-imol-i)* Cp° ( c a l d e g - i m o l - i ) a Electronic spectrum Ground state 2 Π 3 / 2 Microwave spectrum Esr spectrum

»(A)

a>e(cm_1) Dissociation energy, DQ° (kcal mol" l ) determined spectroscopicallyb lonization energy (kcal mol" 1 ) Electron affinity (kcal mol" *) Dipole moment, D Force constant (mdyne Ä - 1 )

ClO

BrO

24-34 23-45 54-14 7-52 b

3006 25-87 56-75 7-67 b

c e l-569±0-001 c [see text]1

d f l-721±0-001 d 713b

63-31 ± 0 0 3 239-88 67-09h l-239 c 6-41*

55-2±0-6 611 1-55* 3-99b

IO

41-84 35-80 58-65 7-86 b

f [r0 = 1-868 ±0-001 f ] 681-4b

42±5

3-87*

a

Selected Values of Chemical Thermodynamic Properties, N . B . S . Technical N o t e 270-3 (1968). R. A . Durie and D . A . Ramsey, Canad. J. Phys. 36 (1958) 4 5 ; R. A . Durie, F . Legay and D . A . Ramsey, ibid. 38 (1960) 4 4 4 . c T. A m a n o , S. Saito, E . Hirota, Y . Morino, D . R. Johnson and F . X . Powell, / . Mol. Spectroscopy, 3 0 (1969) 275. d F . X . Powell and D . R. Johnson, / . Chem. Phys. 50 (1969) 4596. e A . Carrington and D . H . Levy, / . Chem. Phys. 4 4 (1966) 1298. f A . Carrington, P. N . Dyer and D . H . Levy, / . Chem. Phys. 52 (1970) 309. * Measured by mass spectroscopy: V. H . Dibeler, R. M . Reese and D . E. Mann, / . Chem. Phys. 27 (1957) 176. h Calculated from the heat o f hydration of OC1": V. I. Vedeneyev, L . V. Gurvich, V. N . Kondrat'yev, V. A . Medvedev and Y e . L. Frankevich, Bond Energies, lonization Potentials and Electron Affinities, Edward Arnold, London (1966). 1 L . Andrews and J. I. Raymond, / . Chem. Phys. 55 (1971) 3087, 1 G. V. Buxton and F . S. Dainton, Proc. Roy. Soc. A304 (1968) 427. b

contrast with 0 2 F 2 , which is noted for its strong O-O bond, (C10)2 is but a weakly associated species, resembling (NO)2 578. A molecule of composition C1202 (probably Cl-Cl^ ) ^O is postulated as an intermediate in the reaction of C102 - with chlorine or hypochlorites. BrO has been detected by its ultraviolet-visible absorption spectrum following the flash photolysis628 or pulse radiolysis615 of aqueous bromite or hypobromite solutions: similar treatment of I0 2 ~ is impossible, but IO is formed in the reactions and

IO2+IO3-->IO+IO4Ϊ2+ΙΟ-

->IO+2I"

which occur after the flash photolysis or pulse radiolysis of, respectively, aqueous iodate and aqueous hypoiodite solutions629. The spectra and reactions of both radicals have been 628 o . Amichai, G. Czapski and A . Treinin, IsraelJ. Chem. 7 (1969) 351. 629 o . Amichai and A . Treinin, / . Phys. Chem. 7 4 (1970) 8 3 0 .

THE OXIDES OF THE HALOGENS

1385

studied in solution. BrO oxidizes bromite BrO+Br02 ~ -> BrO - + Br02 and disproportionates during hydrolysis 615 : 2BrO+H20 -> BrO + Br02 ~ + 2Η +

ClOO The peroxychlorine radical is an important species in the chain reactions of oxychlorine compounds, e.g.

cio+cio->cioo+ci and as such has been identified by its uv spectrum, using the technique of "modulation kinetic spectroscopy"630. Since its dissociation into chlorine atoms and oxygen molecules is slightly endothermic ( Δ # = 8 ± 2 kcal mol - 1 ) 6 3 1 , ClOO is sufficiently long-lived to be a very useful intermediate; for example, it is responsible for the high efficiency shown by oxygen in promoting the recombination of chlorine atoms. Cl+0 2 +M-*C100+M ClOO+Cl ->Cl 2 +0 2

The heat of formation of the gaseous molecule is estimated to be 21 ±2 kcal mol"1; ClOO is therefore somewhat less endothermic than the more familiar isomer chlorine dioxide. ClOO is also formed when chlorine dioxide is photolysed at low temperature in rigid inert matrices591. hv C10 2 -*[C10+0]->C100 cage

Esr and infrared studies show the isolated molecule (Table 54) closely to resemble OOF: the odd electron occupies a π* orbital; the O-O bond is strong, and the O-Cl bond weak. The formation of chlorine dioxide in the y-radiolysis of KC104 involves ClOO as an inter­ mediate610. Br02andI02 The radicals X 0 2 are produced in aqueous solution (i) by flash photolysis of X 0 3 ~ (X = Cl,BrorI) 6 32 : X0 3 ,H 2 0 -* [X03-,H20]* -* X02+OH+OH(ii) by y-radiolysis of X 0 3 ~ (X = Br or I)6i5.629 Br0 3 - +e"(aq) -► [Br032"] -> Br0 2 +02"

Br0 2 and I 0 2 may also be formed in secondary processes following photolysis or radiolysis of other anions. The radicals have been characterized by their ultraviolet-visible absorption spectra (Table 54). Whereas C102 may be recovered unchanged from aqueous solution, Br0 2 and I 0 2 have very short lifetimes. The rate of hydrolysis of Br0 2 has been explored (see p. 1377), while it has been reported that I0 2 oxidizes iodate to periodate, itself being reduced to IO (see above)629. 630 H. S. Johnston, E. D. Morris, jun. and J. Van den Bogaerde,/. Amer. Chem. Soc. 91 (1969) 7712. «I S. W. Benson and J. H. Buss, /. Chem. Phys. 27 (1957) 1382. 632 F. Barat, L. Gilles, B. Hickel and J. Sutton, Chem. Comm. (1969) 1485.

1386

CHLORINE, BROMINE, IODINE AND ASTATINE: A. J. DOWNS AND C. J. ADAMS

Br03 Ö03

y-Radiolysis of KBr0 3 at 77°K or of frozen alkaline bromate solutions produces a paramagnetic centre identified by its esr spectrum as Br0 3 . The derived O-Br-0 bond angle of 114° indicates considerable flattening of the pyramid relative to the parent anion, and is consistent with trends along the series P0 3 2 - (110°), S0 3 - (111°), C103 (112°) and As0 3 2 (110°), S e 0 3 - (112°), Br0 3 (114°).

3. THE O X Y F L U O R I D E S OF THE H A L O G E N S

The twelve halogen oxyfluorides which have so far been identified contain pentavalent or heptavalent chlorine, bromine or iodine; derivatives of the trivalent halogens have not yet been synthesized. TABLE 56. OXYFLUORIDES OF THE HALOGENS

C102Fal> C10F3 al> C103Fa? C10 2 F 3 a ' c C103OF a b c

Br0 2 F Br0 3 F a ?

I02Fab IOF 3 I0 3 F a ? I02F3b IOF5

Compound functions as donor of F~. Compound functions as acceptor of F~. Two isomers known.

The compounds are generally obtained by fluorination of an appropriate halogen oxide, oxyacid or oxysalt. As might have been expected by interpolation from the corresponding oxides and fluorides, the compounds are oxidizing and fluorinating agents; the chlorine oxyfluorides in particular have been extensively investigated as potential oxidants for rocket fuels. They also tend to function as donors or acceptors of F ~ in forming complexes with, respectively, Lewis acids and fluoride ion-donors, and hydrolyse to an oxyacid maintaining the oxidation state of the halogen. The literature up to 1962 has been reviewed575a, while the oxyfluorides of chlorine have been the subject of a more recent, comprehensive survey572. Definitive structural characterization is restricted to gaseous C103F and the solids IOF 3 and KI0 2 F 2 . However, vibrational spectra and valence force constants have been deter­ mined for many chlorine oxyfluorides, for the most part by Christe and his coworkers; chlorine-oxygen and chlorine-fluorine stretching force constants are listed in Table 57. Both neutral and charged species contain strong Cl-0 bonds (force constants ^ that of CIO4 -), but considerable weakening of the Cl-F bonds occurs with coordination by F -. Halogenyl Fluorides Preparation The three compounds C102F, Br0 2 F and I0 2 F are obtained by fluorination of an oxide

THE OXYFLUORIDES OF THE HALOGENS

1387

TABLE 57. VALENCE FORCE CONSTANTS OF SOME CHLORINE OXYFLUORIDES AND RELATED SPECIES

Molecule

/r(Cl-O) (mdyneÄ - 1 )

/r(Cl-F) (mdyne A - 1 )

6-85

2-59

C10F 2 +

11-21

3-44

CIOF3

9-37

CIOF4-

913

CKV

8-96

C10 2 F

9 07

2-53

C10 2 F 2 CIO3F

8-3 9-41

1-6 3*91

ClOF

C1F2 + CIF3 CIF4CIF5

3 16(eq) 2-34(ax) 1-79

4-74 4-2(eq) 2-7(ax) 211 3-47(ap) 2-67(bas)

Reference L. Andrews, F. K. Chi and A. Arkell, /. Amer. Chem. Soc. 96 (1974) 1997. K. O. Christe, E. C. Curtis and C. J. Schack, Inorg.Chem. 11 (1972) 2212. K. O. Christe and E. C. Curtis, Inorg. Chem. 11 (1972) 2196. K. O. Christe and E. C. Curtis, Inorg. Chem. 11 (1972) 2209. K. O. Christe, C. J. Schack, D. Pilipovich and W. Sawodny, Inorg. Chem. 8 (1969) 2489. D. F. Smith, G. M. Begun and W. H. Fletcher, Spectrochim. Acta, 20 (1964) 1763. K. O. Christe and E. C. Curtis, Inorg. Chem. 11 (1972) 35. W. Sawodny, A. Fadini and K. Ballein, Spectrochim. Acta, 21 (1965) 995. K. O. Christe and C. J. Schack, Inorg. Chem. 9 (1970) 2296. R. A. Frey, R. L. Redington and A. L. K. Aljibury, / . Chem. Phys. 54 (1971) 344. K. O. Christe and W. Sawodny, Z. anorg. Chem. 374 (1970) 306. K. O. Christe, E. C. Curtis, C. J. Schack and D. Pilipovich, Inorg. Chem. 11 (1972) 1679.

or oxysalt of the tetravalent or pentavalent halogen. The most convenient syntheses are C10 2 +AgF 2 KBr0 3 + BrF 5 I2O5 + F2

20°C

-*C10 2 F+AgF

-50°C -> B r 0 2 F + K B r F 4 + i 0 2 20°C

-*2I02F+£02 liqHF

but in each case other oxyhalogen substrates and/or fluorinating agents may be used575»581. Synthesis of C102F by fluorination of C120 or C1206 probably involves initial conversion to C102. I0 2 F is formed by dismutation of IOF 3 . no°c 21OF3

>I02F+IF5

Properties Table 58 lists some properties of the halogenyl fluorides. The vibrational spectra of chloryl fluoride confirm the presence of discrete ¥C\(

molecules of Cs symmetry in the X) vapour and liquid; I0 2 F is likewise revealed as a polymeric solid. Reactions Thermolysis of C102F in a Monel system is appreciable only above 250°C, producing CIF and oxygen633. Bromyl fluoride, however, is unstable above its melting point, and decomposes vigorously at 56°C, probably according to the equation 3Br0 2 F -> B r F 3 + B r 2 + 3 0 2

1388

CHLORINE, BROMINE, IODINE AND ASTATINE: A. J. DOWNS AND C. J. ADAMS TABLE 58. PROPERTIES OF THE HALOGENYL FLUORIDES

Property Colour Melting point (°C) Boiling point (°C) Decomposition temp. (°C) Vapour pressure, log p(mm) = Trouton's constant ( c a l d e g - 1 mol" 1 ) Ai^vapCkcalmol"1) Infrared spectrum Raman spectrum Force constants (mdyne Ä - 1) /r(Cl-O) /r(Cl-F)

C10 2 F

Br0 2 F

I02F

colourless -115a

yellow

colourless >200b (decomposition)

_ 9 b

-6a >250c 8-23-1412/r a

56 b

ca. 200 b

23-2 a 6-2a gas f liq. f

solidd-e solide

907f 2-53f

a

H. Schmitz and H. J. Schumacher, Z. anorg. chem. 249 (1942) 238. M. Schmeisser and K. Brändle, Adv. Inorg. Chem. Radiochem. 5 (1963) 41. c Y. Macheteau and J. Gillardeau, Bull. Soc. chim. France (1969) 1819. d P. W. Schenk and D. Gerlatzek, Z. Chem. 10 (1970) 153. 0 H. A. Carter and F. Aubke, Inorg. Chem. 10 (1971) 2296. f D. F. Smith, G. M. Begun and W. H. Fletcher, Spectrochim. Ada, 20 (1964) 1763. b

Both compounds attack glass at room temperature, the former slowly, the latter rapidly. Hydrolysis of C102F, Br0 2 F or I0 2 F produces the appropriate halate ion: XO2F+2OH- - + X O 3 - + F - + H 2 O

Hydrolysis of Br0 2 F can proceed with explosive violence, while in the hydrolysis of chloryl fluoride ClO? has been detected as an intermediate. ClOF

[cio 2 ] + [sbh 6 ]-

[CIOj + [AsF]-_

AsR

[cio 2 ] + [PtF 6 ]"+[cio 2 p;] + [PtF 6 ]

CKXF

[CIOjt[SnF 6 J

CsF 80°C

Cs[ClQ,Fj

[ClOj [ M F n + 1 ] " A1F

(n = 3, M = B; n = 5, M = P, As, Sb, V)

SCHEME 6. Reactions of chloryl fluoride. 633 Y . Macheteau and J. Gillardeau, Bull. Soc. chim. France (1969) 1819.

THE OXYFLUORIDES OF THE HALOGENS

1389

The stabilities of the complexes formed by CIO2F with Lewis acids parallel the strengths of the acceptors, viz. SbF5 > AsF5 > PF 5 > BF 3 > VF 5 ; on the basis of their vibrational spectra, the adducts are formulated as salts containing discrete [C10 2 ] + cations. The heat of dissociation of the BF 3 complex, that is, for the reaction [C102] + [BF 4 ]-(s) -> C10 2 F(g)+BF 3 (g)

is 24 kcal mol - 1 634. Chloryl fluoro-complexes are commonly products of oxidation or fluorination reactions effected by C102F (Scheme 6). While bromyl fluoride apparently does not react with BF3, AsF5 or SbF5, iodyl fluoride combines with acceptor molecules in the presence of a suitable solvent. [I0 2 ] + [SO3F]" «

SO3

reflux

I02F

AsF 5 liq HF

> [I0 2 ] + [AsF6] -

The structural chemistry of chloryl and iodyl salts was reviewed in Section 4A (p. 1352). C10 2 F and PtF 6 combine together at 25° in two competing reactions^: and

2C10 2 F+2PtF 6 -> [C102] + [PtF 6 ]- + [C10 2 F 2 ] + [PtF6]" 2C10 2 F+2PtF 6 -> 2[C10 2 ] + [PtF6]" + F 2

The chlorine(VII) salt comprises about 10% of the solid product. Acceptor properties are evident in the formation of Cs[C102F2]636 (from CsF and C102F) and KPO2F2]637 (from KF and IO2F in anhydrous HF). The spectroscopically characterized C10 2 F 2 " anion638 has C2v symmetry, the framework being derived from a trigonal-bipyramid with two weak axial Cl-F bonds (force constant 1 -6 mdyne Ä _1 ) and two strong equatorial Cl-O bonds (force constant 8-3 mdyne A - 1 ) ; the remaining equatorial position is occupied by the lone pair of electrons. A similar IO2F2" anionic unit has been defined by X-ray crystallography639 and infrared and Raman spectroscopy545 in KI0 2 F 2 [Fig. 32(a)].

(a) (b) FIG. 32. Molecular structures of (a) I0 2 F 2 ~ (in KI0 2 F 2 ) and (b) IOF 3 . 634 K. O. Christe, C. J. Schack, D. Pilipovich and W. Sawodny, Inorg. Chem. 8 (1969) 2489. 635 K. O. Christe, Inorg. Nuclear Chem. Letters, 8 (1972) 453. 636 D. K. Huggins and W. B. Fox, Inorg. Nuclear Chem. Letters, 6 (1970) 337. 637 j . j . Pitts, S. Kongpricha and A. W. Jache, Inorg. Chem. 4 (1965) 257. 638 K. O. Christe and E. C. Curtis, Inorg. Chem. 11 (1972) 35. 639 L. Helmholz and M. T. Rogers, / . Amer. Chem. Soc. 62 (1940) 1537.

1390

CHLORINE, BROMINE, IODINE AND ASTATINE! A. J. DOWNS AND C. J. ADAMS

Halogen Oxide Trifluorides CIOF3640'641 is prepared (i) by direct low-temperature fluorination of covalent inorganic hypochlorites such as C120 or C10N0 2 , or (ii) photochemically from gaseous mixtures, e.g. CIO3F/F2, CI2/O2/F2, CIF/IOF5 or C102F/CIF5. The pure compound is a colourless gas at room temperature, is an excellent oxidizing fluorinating agent, and is allegedly more corrosive than C1F3. Measured physical properties are: m.p., -42°C; b.p., 29°C; vapour pressure, log p(mm) = 8-433- 1680/Γ; A// v a p = 7-7 kcal mol - 1 ; Trouton's constant = 25-4 cal deg - 1 mol - 1 . Defined spectroscopically in the vapour phase642, the CIOF3 molecule has a Cs structure analogous to that of IOF3; the Cl-0 bond is equatorial and fairly strong (force constant 9-37 mdyne A" 1 ). The 19 F nmr spectrum of the vapour features a single line at — 327 ppm relative to CCI3F (external), which shifts to —272 ppm and broadens for the neat liquid; no splitting is observed on cooling the sample. Molecular association via the axial fluorine atoms is indicated for the condensed phases. The oxyfluoride reacts readily with the Lewis acids BF3, AsF5 and SbF5 to give com­ plexes containing the C10F 2 + cation, structurally akin to SOF2643, and with alkali-metal fluorides forms salts M +[C10F4] ~ (M = K or Cs) in which the anion has C4v symmetry644. Iodine pentoxide dissolves in boiling IF 5 ; white hygroscopic needles of IOF3 separate on cooling the solution. I205 + 3IF5^5IOF3

The crystal structure of the solid contains molecular species (Fig. 32) linked by weak I-F-I bridges645. Vibrational spectra have also been recorded646. On warming to 110°C, IOF 3 dismutates reversibly into IF 5 and iodyl fluoride. 2IOF3->IF5 + I02F

Perhalogenyl Fluorides The compounds CIO3F and IO3F were first recognized in the early 1950s; not synthesized until after the discovery of perbromates (p. 1451), Br0 3 F still awaits detailed investigation. Synthesis The most convenient syntheses involve fluorination of the perhalic acid or its salts. KCIO4

S b F 5 o r HSO3F

> CIO3F

SbF5

KBr04

► Br03F F2

HIO4 liquid H F

>I03F

The reactions possibly proceed via the intermediate formation of the XO3 + cation (X = Cl, Br or I). 640 R . B o u g o n , J. Isabey a n d P . Plurien, Compt. rend. 271C (1970) 1366. 641 D . Pilipovich, C . B . L i n d a h l , C . J . Schack, R . D . Wilson a n d K . O . Christe, Inorg. Chem. 11 (1972) 2 1 8 9 ; D . Pilipovich, H . H . Rogers a n d R . D . Wilson, ibid. 11 (1972) 2 1 9 2 . 642 K . O . Christe a n d E . C . C u r t i s , Inorg. Chem. 11 (1972) 2196. 643 K . O . Christe, E . C . Curtis a n d C . J . Schack, Inorg. Chem. 11 (1972) 2212. 644 K . O . Christe a n d E . C . C u r t i s , Inorg. Chem. 11 (1972) 2209. 645 J . w . Viers a n d H . W . Baird, Chem. Comm. (1967) 1093. 646 H . A . C a r t e r a n d F . A u b k e , Inorg, Chem. 10 (1971) 2296.

THE OXYFLUORIDES OF THE HALOGENS

1391

Perchloryl Fluoride Properties The interesting and useful compound CIO3F has been the subject of three comprehen­ sive reviews572»647'648, and its physical properties have been measured in some detail (Table 59). Electron diffraction has defined the C3v molecule in the gas phase (Fig. 33), while microwave measurements imply a dipole moment of no more than 0Ό23 D 649 . The symmetry of the electric field experienced by the central chlorine atom is also demonstrated by the measured 35G1 and 37C1 nuclear quadrupole coupling constants and by the 35Q-F and 37C1-F spin-spin coupling observed in the 19 F nmr spectrum. The 19 F chemical shift of CIO3F (-287 ppm relative to CC13F) compares with that of molecular fluorine (-428-7 ppm relative to CCI3F). CIO3F offers the highest resistance to electrical breakdown known for any gas, and has been used as an insulator in high-voltage systems. TABLE 59. PROPERTIES OF PERCHLORYL FLUORIDE, CIO3F

Physical properties and thermodynamic parameters Melting point (°C) -147-74» Mff° at 25°C (kcal m o l ' i ) -5-7c Boiling point (°C) -46-67» AG/° at 25°C (kcal mol"i) 11 -5C r c r l t (°C) 95-17* 5°at25°C(caldeg-imori) 66-65c PcritCatm) 530b Cp° at 25°C(caldeg-i mol"*) 15-52c A/fusion (kcal mol-i) 0-9163» A# v a p at b.pt. (kcal mol"i) 4-619» Trouton's constant (caldeg-imol-i) 20-4» Vapour pressure equation:» logp(mm) = - 1652-37/Γ-8-62625 log Γ+00046098Γ+29-44780 for Γ = 164-229°K. Liquid density, /> (gcm~ 3 ): d p = 2-266-1-603Χ10-3Γ-4080Χ10-6Γ2for J = 131-234°K. Viscosity, η (centipoise) : e log η = 299/Γ-1-755 for T = 196-327°K. Surface tension (dyne cm~i): e 24-1-21-3 for T= 198-218°K. Molecular spectra, etc. Infrared: gasf Raman: gas* liquid' i9Fnmr: liquid11 Mass spectrum: gas1 Microwave spectrum: gasJ Electron diffraction: gas k Molecular parameters Bond dissociation energies (kcal mol" 1 )· 1 Z)(Cl-0), 57; D(C\-F), 60 Vibrational frequencies (cm~i):* 1062 niß) 1314 n (01) 716 573 V2 (<*l) "5W 549 414 V3 (öl) "6 W 1 1 Valence force constants (mdyne Ä" ): /r(Cl-O) 9-41; /r(Cl-F) 3-91 Bond lengths (Ä):k«m r f (l)(Cl-0) 1-404 ± 0 0 0 2 r,(l)(Cl-F) l-619±0-004 647 j . F . Gall, Kirk-Othmer Encyclopedia of Chemical Technology, 2nd edn., Vol. 9, p. 598, Interscience, New York (1966). 648 v . M. Khutoretskii, L. V. Okhlobystina and A. A. Fainzil'berg, Russ. Chem. Rev. 36 (1967) 145. 649 A. A. Maryott and S. J. Kryder, / . Chem. Phys. 27 (1957) 1221; D . R. Lide, jun. ibid. 43 (1965) 3767.

1392

CHLORINE, BROMINE, IODINE AND ASTATINE: A. J. DOWNS AND C. J. ADAMS

Table 59 (cont.)

19

F chemical shift (ppm relative to CCI3F as internal standard) : h 287 /(Cl-F)(Hz): h 278 ± 5 Nuclear quadrupole coupling constants, e2Qq (MHz): J 35C1, -19·2±0·5; Ionization potential (kcal mol" 1 ): 1 314 ± 4 Dipole moment, D:n 0023 ± 0 0 0 3

37Q,

-15·4±1·5

a

J. K. Koehler and W. F. Giauque, / . Amer. Chem. Soc. 80 (1958) 2659. A. Engelbrecht and H. Atzwanger, / . Inorg. Nuclear Chem. 2 (1956) 348. c Selected Values of Chemical Thermodynamic Properties, N.B.S. Technical Note 270-3 (1968). d R. L. Jarry, / . Phys. Chem. 61 (1957) 498. e J. Simkin and R. L. Jarry, / . Phys. Chem. 61 (1957) 503. f D. R. Lide, jun., and D. E. Mann, / . Chem. Phys. 25 (1956) 1128; F. X. Powell and E. R. Lippincott, ibid. 32 (1960) 1883. β H. H. Claassen and E. H. Appelman, Inorg. Chem. 9 (1970) 622. 11 S. Brownstein, Canad. J. Chem. 38 (1960) 1597; H. Agahigian, A. P. Gray and G. D . Vickers, ibid. 40 (1962) 157; J. Bacon, R. J. Gillespie and J. W. Quail, ibid. 41 (1963) 3063; A. A. Maryott, T. C. Farrar and M. S. Malmberg, / . Chem. Phys. 54 (1971) 64. 1 V. H. Dibeler, R. M. Reese and D. E. Mann, / . Chem. Phys. 11 (1957) 176. J D. R. Lide, jun., / . Chem. Phys. 43 (1965) 3767. k A. H. Clark, B. Beagley, D. W. J. Cruickshank and T. G. Hewitt, J. Chem. Soc. (A) (1970) 872. 1 W. Sawodny, A. Fadini and K. Ballein, Spectrochim. Acta, 21 (1965) 995. m For bond angles see Fig. 33. n A. A. Maryott and S. J. Kryder, / . Chem. Phys. 27 (1957) 1221. b

Θ

141(1) Ä

1-619(4) A

(a)

(b)

FIG. 33. Molecular structures of (a) CIO3F and (b) [NC10 3 ]2- (in K2NCIO3).

Reactions Perchloryl fluoride is often regarded as an inert substance, although in fact many of its reactions occur under mild conditions. While AG/° is slightly positive, Δ///° is negative, so that spontaneous decomposition at room temperature is virtually impossible, and the compound is stable towards explosion; thermal decomposition is noticeable only above 400°C. The oxidizing powers of CIO3F depend strongly on the nature of the reaction

1393

THE OXYFLUORIDES OF THE HALOGENS

medium and on the temperature. In 0-1 M acid, iodide ion is quantitatively oxidized to iodine. CIO3F+8I- + 6 H + -> C\~ + F - + 4 I 2 + 3 H 2 0

Many materials towards which it is inert at 20°C are rapidly attacked at 150-200°C. Perchloryl fluoride is used in large quantities, either alone or admixed with halogen fluorides, as an oxidant for rocket fuels; mixtures of CIO3F with C1F5 yield CIOF3 on ultraviolet photolysis641. CIO3F is highly susceptible to nucleophilic attack at the chlorine atom. Under aqueous conditions hydrolysis occurs only on heating the compound to 200°C in a sealed tube with concentrated alkali, but in alcoholic potash hydrolysis is rapid and quantitative at room temperature. CIO3F reacts smoothly with liquid or aqueous ammonia: CIO3F + 3NH3 -> NH 4 [NHC10 3 ] + NH 4 F

From aqueous solutions of this ammonium salt other amidoperchlorates may be prepared, e.g. MINHC103, M^NCIC^ and M n NC10 3 (M1 = Na, K, Cs or Ag; M11 = Sr,Ba or Pb); many of the solids are explosive. Structural investigations650 of these amidoperchlorates include an X-ray study of crystalline K 2 NC10 3 ; the salt is isomorphous with K 2 S0 4 , and the anion approximates to C^v symmetry (Fig. 33)651. Although the unstable explosive liquid N-perchloropiperidine has been prepared, other organic amines undergo oxidative decom­ position with CIO3F. In organic chemistry CIO3F has been widely used as a mild fluorinating agent. While phenyllithium displaces F - from CIO3F in a straightforward manner (Scheme 7), reaction with carbanions involves C-F bond-formation in many cases where the anion is stabilized by resonance with one or more neighbouring electronegative centres, as in diesters and polynitro compounds. R 3 C- + CIO3F -> R3CF+CIO3-

The probable mechanism involves attack on CIO3F by the more nucleophilic of the anionic sites, followed by rapid formation of the strong C-F bond652. R Et02C-C=C-OEt

R Et02OC^C-OEt

R Et0 2 OC-C-OEt

df$


<φ>-

A related mechanism is proposed for the oxofluorination of carbon-carbon bonds. Unlike C102F, perchloryl fluoride does not form complexes with Lewis acids such as SbF5 or S0 3 ; neither is it soluble in liquid HF. In the presence of Friedel-Crafts catalysts, however, CIO3F inserts the CIO3 group into aromatic nuclei (p. 1573), a reaction which may involve the intermediate agency of CIO3+. Perbromyl Fluoride653 Br0 3 F is a colourless reactive gas, solidifying at about -110°C to a white solid. At 650 j . Goubeau, E. Kilcioglu and E. Jacob, Z. anorg. Chem. 357 (1968) 190; A. I. Karelin, Yu. Ya. Kharitonov and V. Ya. Rosolovskii, Zhur. priklad. Spektroskopii, 8 (1968) 256, 458; Russ. J. Inorg. Chem. 13(1968) 1234. 65i N. I. Golovina, G. A. Klitskaya and L. O. Atovmyan, / . Struct. Chem. 9 (1968) 817. 652 w . A. Sheppard, Tetrahedron Letters (1969) 83.

1394

CHLORINE, BROMINE, IODINE AND ASTATINE: A. J. DOWNS AND C. J. ADAMS

room temperature it slowly decomposes in metal or Kel-F containers. As with CIO3F, the infrared and Raman spectra of the gaseous compound imply a molecule of C3v symmetry654. It is hydrolysed by aqueous alkali. [NH] + [NHC10jC10>F

CtOF

^

4

3

C/f,

MeCOCF CO Et

SCHEME 7. Reactions of perchloryl

fluoride.

Periodyl Fluoride575a IO3F is a white crystalline solid, stable in glass, and decomposing at 100°C to I0 2 F and oxygen. It is soluble in liquid HF, reacting therein with BF 3 and AsF 5 to give materials which may be formulated as [I03][BF4] and [IO3][AsF6],10HAsF6, respectively. SO3 reduces IO3F to iodyl fluorosulphate: IO3F+SO3 -> I O 2 S O 2 F + 0 2

Like CIO3F, periodyl fluoride reacts with ammonia, but the product, NH3,I03NH2, has not been characterized. Iodine Oxypentafluoride Iodine oxypentafluoride, IOF 5 , is obtained by the reaction of IF 7 with water, silica, glass or I 2 0 5 6 5 5 ; though not susceptible to hydrolysis, the colourless compound, m.p. 4-5°C, is difficult to purify. Spectroscopic investigations are consistent with a structure of C^ symmetry (6)656. 653 E. H . Appelman and M. H . Studier, / . Amer. Chem. Soc. 91 (1969) 4561. 654 H . H . Claassen and E. H . Appelman, Inorg. Chem. 9 (1970) 622. 655 R. j . Gillespie and J. W. Quail, Proc. Chem. Soc. (1963) 278; N . Bartlett and L. E. Levchuk, ibid. p. 3 4 2 ; J. H . Holloway, H . Selig and H . H . Claassen, J. Chem. Phys. 54 (1971) 4305. 656 D . F. Smith and G. M. Begun, / . Chem. Phys. 4 3 (1965) 2001.

THE OXYFLUORIDES OF THE HALOGENS

ν

Λ

1395

F

F (6)

Normal coordinate analysis of the vibrational frequencies implies little difference between the axial and equatorial I-F bonds (force constants 4-60 and 442 mdyne A - 1 , respectively), though the 19 F chemical shifts clearly discriminate between axial and equatorial fluorine atoms (relative to SiF4, the values are — 272 and — 236 ppm, respectively)657. The microwave spectrum has not been analysed to yield interatomic distances, but the dipole moment is given as 1-08 ±0-11)658. Halogen Dioxide Trifluorides C102F3 Two materials of composition CIO2F3 have been described in the literature. (1) Treatment of [C10 2 ]+[PtF 6 ]- containing .about 10% [C102F2]+ [PtF6] - (p. 1389) with NOF in a sapphire reactor at — 78°C displaces, among other volatile products, a material of composition CIO2F3, colourless in the solid, liquid and vapour phases, and apparently stable at 25°C<>59. [CIO2F2] + [PtF 6 ]" + NOF -> CIO2F3+[NO] + [PtF 6 ]-

Its infrared spectrum is consistent with a molecule of C2v symmetry (cf. I0 2 F 3 below) having equatorial oxygen atoms. The existence of C10 2 F 2 + indicates fluoride-donor abilities which have yet to be directly tested by experiment. (2) A violet solid 0 2 C1F 3 660, a vigorous oxidizing agent even at low temperature, has been prepared by two routes: hv

CIF3 + O2

195βκ

> O2CIF3 <

CIF + O2F2

120°κ

The compound is formulated on spectroscopic grounds as FOOClF2; with excess C1F, 0 2 F 2 gives a blue compound believed to be F2C100C1F2. I02F3 Iodine dioxide trifluoride has been prepared by the following route661. Ba 3 H 4 (I0 6 )2

HS0 3F

> [HIO2F4]

S0 3

► IO2F3

19

The yellow sublimable solid melts at 41°C; F nmr studies of the melt are consistent with the presence of both Cs (7) and C2v (8) isomers.

F-f«g

F-I^S

O

F

(?)

(8)

657 N . Bartlett, S. Beaton, L. W . Reeves and E . J. Wells, Canad. J. Chem. 4 2 (1964) 2531. 658 s . B . Pierce and C. D . Cornwell, / . Chem. Phys. 47 (1967) 1731. 659 K. O. Christe, Inorg. Nuclear Chem. Letters, 8 (1972) 457. 660 A . V. Grosse and A . G. Streng, Chem. Abs. 66 (1967) 39492z; ibid. 67 (1967) 75007z. 661 A . Engelbrecht and P. Peterfy, Angew. Chem., Internat. Edn. 8 (1969) 768.

1396

CHLORINE, BROMINE, IODINE AND ASTATINE: A. J. DOWNS AND C. J. ADAMS

Like IOF 5 , the compound is quite resistant to hydrolysis. On exposure to sunlight it liberates oxygen (and some ozone). I02F34lOF3 + i02

Fluorine Perchlorate While fluorine perchlorate has the empirical composition of a halogen oxyfluoride, it has no halogen-fluorine bond, and is considered below (p. 1473) alongside the other halogen perchlorates. 4. O X Y A C I D S A N D O X Y S A L T S OF THE (A)

HALOGENS

INTRODUCTION

Listed in Table 60 are the known oxyacids of the halogens. Many of these acids and their salts were first discovered, or at least recognized for themselves, during the great advance in synthetic and descriptive inorganic chemistry which marked the period 1770-1830. Only recently has the list been completed, however, when the successful syntheses of hypofluorous acid662, perbromic acid663 and perbromate salts663 exorcised a considerable accumulation of conjecture about their "non-existence". Many of the acids are known only in solution (Table 60), but, except for iodites and hypoiodites, reasonably stable crystalline salts have been isolated; the existence of HIO2 and IO2 ~ as more than transient species in solution is doubtful. The molecules HOX (X = F, Cl or Br), HIO3 and HXO4 (X = Cl, Br or I) are sufficiently stable in the vapour phase to permit their investigation by mass spectroscopy and other techniques. A series of hydrates of perchloric acid has been characterized, while iodic acid forms a 1:1 compound with iodine pentoxide, HIO3J2O5. TABLE 60. T H E OXYACIDS OF THE HALOGENS

HOF f t

a

b c d e

HOClbe HCIO2 c HCIO3 e HCIO4 b ' c - d ' e

HOBrb^ H B r 0 2 ( ?) e HBr03 e H B r 0 4 b-e

HOIe H I 0 2 ( ? ?) e HIO3 b « d ' e ΗΙΟ4 b « d Η7Ι3Οΐ4α H 5 I 0 6 d .e

Stable in solid and vapour phases at low temperature; rapidly oxidizes water. K n o w n in vapour phase. K n o w n as pure liquid. K n o w n in solid phase. K n o w n in aqueous solution.

The anhydrous molecular acids HmXOn (X = F, Cl, Br or I) all contain O-H rather than X-H bonds, and similar structures are assumed for the undissociated acids in solution. Periodic acid H 5 I0 6 (formally HI0 4 ,2H 2 0) exists as the ortho-acid (HO)5IO in the solid state and solution, although the solid hydrates of perchloric acid consist of CIO4 ~ anions 662 M. H . Studier and E. H . Appelman, / . Amer. Chem. Soc. 93 (1971) 2349. 663 E. H. Appelman,/. Amer. Chem. Soc. 90 (1968) 1900; Accounts Chem. Res. 6 (1973) 113.

OXYACIDS AND OXYSALTS OF THE HALOGENS

1397

balanced by aquated protons. While H 5 I0 6 behaves as a weak tribasic acid (pK1? 3-29; pK2, 8-3; pK 3 , 11-6), the strengths of the monobasic acids rise as the oxidation state of the halogen rises; the chlorine acids in particular closely follow Pauling's rules664, as shown by the approximate pK a s: HOCl, 7-52; HC102, 1-94; HC103, - 3 ; HC104, - 1 0 . ThepK a s of related acids increase in the sequence: Cl < Br < I. Iodic acid and (especially) periodic acid and their derivatives undergo polymerization in solution, and salts containing oligomeric anions have been prepared. The halogen-oxygen bond lengths in the acid H m XO n or anion XOn ~ contract as the oxidation state of the halogen rises, and the valence force constant undergoes a simultaneous increase. These effects, discussed in Section 4B1, may be explained either in terms of dn-pn bonding in the X-O bond, or by an increasing charge separation Xs+-Os~ in the bond. The charges on the atoms in the anions have been investigated by calculation and by spectroscopic methods (Table 44), but without producing consistent results. Another noteworthy aspect of the structural chemistry of the oxyacids and their deriva­ tives is that higher coordination numbers are found for iodine than for chlorine or bromine; this must be due, at least in part, to the greater size of the iodine atom. While chlorine and bromine are restricted to coordination numbers ^ 4 in the oxyacids and their salts, many periodates contain six iodine-oxygen bonds, and the crystal structures ofiodates (including HIO3) display significant iodine-oxygen contacts at less than the sum of the van der Waals* radii (3-55 A), which increase the number of neighbours to 6, 7 or 8. The ease with which iodine expands its coordination shell has been used to rationalize the observation that redox and exchange reactions involving iodates and periodates occur much more quickly than those of their chlorine and bromine relatives. The readiness with which iodates and peri­ odates ligate metal ions (forming iodato- and periodato-complexes) probably reflects a fairly high charge separation in the I-O bond, producing polarizable oxygen atoms. While chlorates, bromates, (presumably) perbromates and (especially) perchlorates are weaker complexing agents than iodine anions (CIO4 " being so weak that perchlorates are commonly used to establish media of constant ionic strength in studies of the complexation equilibria of other species), it is clear from spectroscopic and crystallographic data that these anions (even C104 ~) will attach themselves to cations if stronger donors are absent. Both thermodynamic and kinetic factors are important in the chemistry of the halogen oxyacids, and particularly in the interrelationships of the various oxidation states of the halogens in the condensed phases; moreover, both the thermodynamic and kinetic param­ eters which control the reactions of the acids and anions in solution are critically sensitive to pH. The examples in the subsequent paragraphs illustrate the importance of thermo­ dynamic, kinetic and environmental features in discussing the chemistry of these systems; many additional examples will be found in the remainder of the section. The thermodynamic properties of oxyacids and oxysalts have already been summarized in connection with the oxidation state diagram and redox potentials for halogen derivatives in aqueous solution (Figs. 2 and 14) and similar patterns are encountered for the solid derivatives; some relevant redox potentials are also tabulated below amongst the properties of the various species. In general, the ease of oxidation of the halogens increases in the order F <ξ Br < Cl < I; for the aqueous halate ions AG/° = +4-55 (Br0 3 "), -1-90 (C103") and -30-2 kcal g-ion -1 (I0 3 _ ). The ready oxidation of iodine is expected on the basis of 604 L. Pauting, The Nature of the Chemical Bond, 3rd edn., pp. 324-328, Cornell University Press, Ithaca (1960).

1398

CHLORINE, BROMINE, IODINE AND ASTATINE: A. J. DOWNS AND C. J. ADAMS

established trends in the Periodic Table. The observation that bromine compounds are thermodynamically less stable than analogous chlorine and iodine compounds is an example of the "alternation effect"665, which rationalizes the "anomalous" behaviour of the elements of the third row in terms of the irregular increase in nuclear charge on descending a Group, with particular emphasis on the shielding characteristics of different orbitals and on the resultant variation in promotion energies. Similar effects are seen in the instability of bromine oxides, fluorides and oxyfluorides. Though the oxyacids and their salts are generally stable thermodynamically with respect to formation from the elements (the notable exceptions with AG/° > 0 being the aqueous Br(>3~ and B r 0 4 _ ions), other decomposition processes are possible. The reaction XO n - - > X - + / * / 2 0 2

is more or less exothermic for chlorine and bromine anions, though for kinetic reasons (rupture of the X - 0 bond is necessary) it is observed at room temperature only for the hypohalites. The oxidizing powers of the compounds are also noted in many reactions in solution, the mechanisms usually involving oxygen-atom transfer. In acid solution only HIO3 is stable with respect to liberation of oxygen from water, although the reactions of the potentially unstable species are very slow unless suitably catalysed. The oxidizing powers of these compounds are exploited in the laboratory in numerous preparative and analytical processes, on a large scale in bleaching and sterilizing operations, and in the manufacture of rocket fuels and explosives. The same oxidizing properties also bring risks of fire and explosion in manipulation, especially in the presence of inflammable organic matter; perchlorate explosions have caused injury or death to many workers. The closeness of successive redox potentials accords with the experimental observation that disproportionation reactions are common. For chlorine, bromine and iodine the + 1 and + 3 oxidation states are unstable with respect to formation of — 1 and + 5 (or 4- 7) states, while in the chlorine systems evolution of C102 is a possibility, especially from acid solution. In general, several different disproportionation processes may be open to each species; the choice of reaction then turns on kinetic rather than thermodynamic factors. The evident effect of pH on the redox potentials (Fig. 14) is best translated into chemical terms by an example. The disproportionation of the halogens in hot alkaline solution, 3X2+6OH- ->5X-+X0 3 -+3H 2 0 is used to prepare halates in high yield; the equilibrium constants are very favourable (Cl, 4x 1074; Br, 2 x 1037; I, 1 x 1023), and the reaction proceeds (via X O - a n d X 0 2 " ) to completion in a matter of hours for chlorate, of minutes for iodate. However, in acid solu­ tion the reverse reaction, X0 3 -+5X-+6H + ->3X2+3H20 has favourable equilibrium constants (Cl, 2x 109; Br, 1 x 1038; I, 6x 1044). In the case of bromate and iodate, the reaction is instantaneous and is used in the analysis of these anions; chlorate is reduced to C102 (in addition to chlorine) by Cl ~ in acid solution. For chlorine and bromine the rates of decomposition, exchange and redox reactions666 generally decrease as the oxidation state of the halogen rises, e.g. ClO ~ > C102 ~ > CIO3 ~> CIO4-. Perchlorates and perbromates are noted for being sluggish oxidizing agents at 665 c . S. G . Phillips and R. J. P. Williams, Inorganic Chemistry, Vol. 1, p. 663. Clarendon Press, Oxford (1965); R. S. N y h o l m , Proc. Chem. Soc. (1961) 273. 666 A . G . Sykes, Kinetics of Inorganic Reactions, pp. 199, 269, Pergamon (1966).

OXYACIDS AND OXYSALTS OF THE HALOGENS

1399

room temperature, and exchange oxygen only very slowly with water: the half-life of C104~ in 6 M HCIO4 with respect to oxygen exchange is more than 100 years at 25°C. How­ ever, periodates are vigorous oxidizing agents and trade oxygen rapidly with water. The rates of reactions are usually faster in acid solution than in neutral or alkaline solution, although it may be hard to distinguish acid catalysis from other changes brought about by decreasing the pH—especially changes in the thermodynamic parameters; it has also been shown that in solutions containing both an acid and its anion, the acid is usually more reactive. Transition metal cations catalyse many decomposition reactions of the oxyacids and oxyanions in solution, probably via intermediate complexes which labilize a halogen-oxygen bond. Finally, the rates of reaction increase as the size of the halogen increases. The equilibrium constants for the disproportionation of the hypohalites in alkaline solution 30X- v±X0 3 -+2X26

are all sufficiently large (Cl, 3 x 10 ; Br, 8 x 1014; I, 5 x 1023) to discount thermodynamic influences; while OCl~ persists for days in solution, OBr~ decomposes in hours, and 01 ~ disappears almost instantaneously. Similar observations have been made for the exchange of oxygen between water and XO3 -, which proceeds at a measurable rate for C103" at 100°C in acid solution, for Br0 3 ~ at 30°C in acid solution, and for I0 3 ~ at 20°C in neutral solution.

(B) HYPOHALOUS ACIDS AND HYPOHALITES

Introduction The story of halogen(I) oxyacids576 begins in 1774, when Scheele remarked that chlorine water was able to bleach vegetable colours. That this property might be useful commercially was suggested in 1785 by Berthollet, who also noted that solutions of chlorine in potash lye were more concentrated and more powerful bleaches than aqueous solutions and did not have the deleterious effects on workers and materials caused by excess chlorine; patents for this bleaching process were taken out in 1789. Influenced by the high cost of alkalis, Tennant in 1789 prepared bleaching solutions by dissolving chlorine in aqueous suspensions of lime, strontia or baryta. He subsequently (1798) patented a process for the manufacture of "bleaching powder" by saturating dry calcium hydroxide with chlorine gas. The chemical constitution of these bleaching solutions wasfirsttruly recognized in 1834 by A. J. Balard, who prepared an aqueous solution of hypochlorous acid and isolated the anhydride CI2O; he had earlier (1821) obtained solutions of hypobromous acid by the gradual addition of yellow mercuric oxide to bromine water. Solutions of hypoiodous acid were similarly generated, but not until 1897. HOI is seen in the mass spectrum of H 5 I0 6 . Of the acids, only hypofluorous acid has been obtained in the pure form, milligram amounts being produced in thefluorinationof ice662; it is unstable at ambient temperatures. Solutions of HOC1, HOBr, HOI or their salts are unstable, as are the solid salts themselves. Nevertheless their oxidizing powers have been widely used in commercial and household bleaching and disinfecting operations, as well as in analytical and preparative procedures in the laboratory. The acids and their derivatives (especially esters) are important halogenating agents in organic chemistry. The following reviews of the chemistry of the halogen(I) oxyacids and their derivatives

1400

CHLORINE, BROMINE, IODINE AND ASTATINE: A. J. DOWNS AND C. J. ADAMS

have recently appeared: hypochlorous acid and hypochlorites572'574'5751*'667; hypobromous acid and hypobromites575b·668'669; hypoiodous acid and hypoiodites575b'616. Preparation Hypohalous Acids The most convenient methods for obtaining aqueous solutions of the acids HOC1, HOBr and HOI involve perturbing the equilibria in the disproportionative hydrolysis of the halogens (p. 1188) X2+H2O ^ H + + X - +HOX

(which normally lies well to the left), by removal of halide ion in the form of an insoluble or sparingly dissociated salt. Although many silver(I) and mercury(II) salts have been used for this purpose, Ag 2 0 and HgO are probably the best reagents, affording the metal halide without introducing extraneous anions into solution. The stability of acid solutions depends on the nature of other ions present; for HOC1 and HOBr, some purification may be effected by distillation of the solution under reduced pressure. Concentrated (> 5 M) solutions of hypochlorous acid may be prepared by treating dichlorine monoxide (either the pure liquid or solutions in CC14) with water at 0°C 577 ; on a large scale the acid is produced by passing C120 gas into water574. Hypohalites Disproportionation of the halogens in alkaline solution is thermodynamically more favourable than in neutral media (p. 1191), and is rapid at room temperature; aqueous solutions containing G O - , BrO" or IO~ result from dissolving the halogen in a cold solution or suspension of the appropriate base. X2+2OH- ^ X - + O X - + H 2 0

Made by this method, the solutions are contaminated by halide ion, and further by the subsequent disproportionation of the hypohalites themselves (see below). The presence of excess base stabilizes the solutions to some extent. Methods have been reported for the electrochemical oxidation of halides to hypohalites in cold dilute solution576. Chemical oxidation of halide ions is also possible: in alkaline solution hypochlorites oxidize bromides to hypobromites, and hypoiodites can be generated from iodide with either hypochlorite or hypobromite. Hypochlorite solutions have been prepared by the neutralization of hypochlorous acid or dichlorine monoxide577. Separation of hypochlorites from chlorides may be effected by treatment with an amine base (or an alcohol), which with OC1 ~ generates the chloramine (or hypochlorite ester) (p. 1410); the organic derivative is readily separated, and the hypo­ chlorite released with alkali. The hydrolysis of N-chloro or N-bromo compounds, e.g. "chloramine-T" (Na+[CH3C6H4S02NC1] -), is an expedient method for the in situ genera­ tion of hypohalite ions. The preparation of pure solid hypochlorites is difficult, if not impossible; procedures have been described for obtaining materials with purity < 90%. Hypochlorite solutions are 667 c . C. Addison, Mellows Comprehensive Treatise on Inorganic and Theoretical Chemistry, Supplement II, Part I, pp. 544-569, Longmans, London (1956). 668 B . Cox, Mellofs Comprehensive Treatise on Inorganic and Theoretical Chemistry, Supplement II, Part I, pp. 750-752, Longmans, London (1956). 669 p. j . M . Radford, Bromine and its Compounds (ed. Z. E. Jolles), pp. 154-158, Benn, London (1966).

OXYACIDS AND OXYSALTS OF THE HALOGENS

1401

sufficiently stable in the absence of light to allow evaporation or crystallization at room temperature577 (at higher temperatures disproportionation dominates), but the solids so derived are more or less contaminated with chloride, chlorite and chlorate; hypochlorites of Li, Na, K, Mg, Ca, Sr and Ba have been reported in both hydrated and anhydrous forms667. Basic hypochlorites of the alkaline earths have been synthesized (and manufactured) by the chlorination of metal hydroxides (either the solid or an aqueous suspension); such salts frequently contain large amounts of chloride. Solid yellow sodium and potassium hypobromites [NaOBr,xH 2 0, where x = 5 or 7, and KOBr,3H20] can be crystallized from the supersaturated solutions which result from the addition of bromine to cold (< 0°C) concentrated (^ 40%) aqueous solutions of NaOH or KOH; the compounds are slightly impure, and are unstable above 0°C669. There are no substantiated accounts of the isolation of solid metal hypoiodites. Structure and Properties Microwave studies (Table 61) have confirmed the HOX structures (as distinct from HXO) of the gaseous hypofluorous acid and hypochlorous acid molecules. The bond angle in HOF is smaller than that in either H 2 0 (104-7°) or F 2 0 (103-2°); it is equally consistent with the essential neutrality of the fluorine atom in HOF that the O-F bond is longer and weaker than in F 2 0 (1-412 Ä). By contrast, the angle in HOC1 is only slightly less than that of the water molecule; within the experimental errors, the O-H and Cl-O bond lengths are identical with those of the parent oxides H 2 0 (0-96 Ä) and C120 (1-70 Ä). The O-Cl bond in CH3OCl is slightly shorter (1-67 Ä; Table 64). The stretching force constants of the O-X bonds in HOX (X = Cl or Br) are much greater than those of the oxides X 2 0 (Table 41); the values for OCl~ and OBr~ (Table 62) fall between the two extremes. The hypohalite ions are isoelectronic with the corresponding diatomic halogen fluorides, but have smaller stretching force constants. The ultravioletvisible spectra of the aqueous ions display weak maxima (listed in Table 62) associated with the transitions 3 Π 0 + <- ιΣ +, and stronger absorptions at shorter wavelengths identified with charge-transfer-to-solvent transitions. Dissociation in the 3 Π 0 + excited state [giving X~(!S) and 0( 3 P)] probably accounts for the high photochemical sensitivity of the OX _ anions. The hypohalous acids are weak acids; this factor, combined with their extreme insta­ bility, makes the accurate determination of dissociation constants a difficult undertaking. The values for HOC1 and HOBr given in Table 61 summarize many individual studies listed elsewhere (HOG 667 · 670 and HOBr66^). Because the acids are so weak, solutions of the salts display an alkaline reaction, the equilibrium OX-+H2O ^ H O X + O H -

lying well to the right. Except at high pH, hypohalite solutions contain significant quantities of the free acid. Concentrated solutions of hypochlorous acid contain large amounts of the anhydride, dichlorine monoxide, which can form a separate liquid layer; the solution at the eutectic 670 j . c . Morris, / . Phys. Chem. 70 (1966) 3798. 671 C. H. Secoy and G. H. Cady, / . Amer. Chem. Soc. 62 (1940) 1036.

Reduction potentials, in acid solution0 £ ° ( X 0 3 - /HOX) (volts) £°(HOX/iX 2 ) (volts) £°(HOX/X-) (volts) Ultraviolet-visible spectrum Kinetic studies of decomposition

+ 1-43 + 1-63 + 1-50 P r

colourless -28-9 -19-1 34 ca. 3 x 1 0 - 8 1

HOX in aqueous solution Colour AHf° (undiss.) (kcal m o l ' i ) * AGf° (undiss.) (kcal mol" 1 )* 5° (undiss.) (cal deg~i mol"i) k Dissociation constant, Κα>29&

+ 1-49 + 1-59 + 1-33 q s

pale yellow -27-0 -19-7 34 ca. 5 x 1 0 - 9 1

7-14* 3-59 0-70 0-64

3537-1 1359-0 8860 7-35* 3-86 0-77 0-45 vapour 3

matrix-isolated 6

-20±10

3590 1164 626

-21-5 c 1-693 + 0-005 0-97 ± 0 0 1 103 ±3° vapour®·f matrix-isolated*

HOBr

3581 1239 729

matrix-isolated d - h

b 1-442 ± 0 0 0 1 0-964±0-01 97-2±0-6°

HOCl

Fundamental frequencies (cm - 1 ) 1 vi O-H stretch V2 Bend v3 O-X stretch Force constants /r(0-H)(mdyneA-i) A(0-X)(mdyneÄ-i) fe (mdyne Ä r a d - 1 ) /r0(mdynerad_1) Ultraviolet-visible spectrum

Molecular HOX Atf/XgHkcalmol-i)* Microwave spectrum r e (0~X)(A) r e (0-H)(A) L H-O-X Infrared spectrum

HOF

TABLE 61. PROPERTIES OF THE HYPOHALOUS ACIDS

t

+ 1-14 + 1-45 + 100

-33-0 -23-7 22-8 ca. IX lO-ii m 4-5x10-13°

-21±10

HOI

V. I. Vedeneyev, L. V. Gurvich, V. N. Kondrat'yev, V. A. Medvedev and Ye. L. Frankevich, Bond Energies, lonization Potentials and Electron Affinities, pp. 78, 130, Edward Arnold, London (1966). b Measured for HOF and D O F ; H. Kim, E. F. Pearson and E. H. Appelman, J. Chem. Phys. 56 (1972) 1. c Measured for H035C1, H037C1, D035C1, D 0 3 7 Q ; D. C. Lindsay, D. G. Lister and D . J. Millen, Chem. Comm. (1969) 950. d Measured for HOF; J. A. Goleb, H. H. Claassen, M. H. Studier and E. H. Appelman, Spectrochim. Acta, 28A (1972) 65. θ K. Hedberg and R. M. Badger, / . Chem. Phys. 19 (1951) 508. f High resolution study of Vi for H035C1, H037C1, D035Q, D037Q gives r(O-Cl) = 1 ·689±0·006 Ä, r(O-H) = 0·97 4 ±0·02 Ä, L H-O-Cl = 104·8±5°; R.A. Ashby, / . Mol. Spectroscopy, 23 (1967) 439. e Includes isotopic data (H, D ; ™0, i»0); HOX produced by photolysis of A r - H X - 0 3 mixtures; I. Schwager and A. Arkell, / . Amer. Chem. Soc. 89 (1967) 6006. h Measured for H F · · · H O F ; P. N . Noble and G. C. Pimentel, Spectrochim. Acta, 24A (1968) 797. 1 Data for HOF isolated in N 2 ; HOC1, HOBr isolated in Ar. i W. C. Fergusson, L. Slotin and D. W. G. Style, Trans. Faraday Soc. 32 (1936) 956. k Selected Values of Chemical Thermodynamic Properties, N.B.S. Technical Note 270-3 (1968). 1 See text. m Determined from kinetic studies; Y. T. Chia, U.S. Atomic Energy Commission Rept. UCRL-8311 (1958); Chem. Abs. 53 (1958) 2914e. n Measured analytically; A. Skrabal, Z. Elektrochem. 28 (1922) 57; Ber. 75B (1942) 1570. ° A. G. Sharpe, Halogen Chemistry (ed. V. Gutmann), Vol. 1, p. 13, Academic Press (1967). p Maxima at 235 nm (e ~ 100 M"i) and 290 nm (c ~ 27 M"i); J. C. Morris, / . Phys. Chem. 70 (1966) 3798. q L. Farkas and F. S. Klein, / . Chem. Phys. 16 (1948) 886. ^ r g M . W . Lister, Canad. J. Chem. 3 0 (1953) 879. 8 <-*> P. J. M. Radford, Bromine and its Compounds (ed. Z. E. Jolles), p. 154, Benn, London (1966). 1 M. L. Josien and C. Courtial, Bull. Soc. chim. France (1949) 374.

a

1404

CHLORINE, BROMINE, IODINE AND ASTATINE! A. J. DOWNS AND C. J. ADAMS TABLE 62. PROPERTIES OF HYPOHALITE IONS IN AQUEOUS SOLUTION

Thermodynamic functions A#/°(kcalmol-i)a &Gf° (kcal mol" 1 )* ^(caldeg-imor1)* Reduction potentials, in basic solution13 £°(X0 3 -/OX-)(volts) E 0 (OX-/*X 2 )(volts) £°(OX-/X-)(volts) Raman frequency (cm - 1 ) Force constant (mdyne Ä - J ) Ultraviolet-visible spectrum Amax(nm) e»«ax(M-l)

Disproportionation, 3OX ~ ^ 2X ~ + XO3 ~, in basic solution Kh Rate

Kinetic studies

oci-

OBr-

OI-

-25-6 -8-8 10

-22-5 -80 10

-25-7 -9-2 -1-3

+0-50 + 0-40 +0-89 713° 3-29c

+ 0-54 + 0-45 + 0-76 620 d 3 0d

292e 350

331 f 326

3 X 1026

Slow at 20°C; fairly rapid at 75°C i

+ 014 + 0-45 + 0-49

365« 31

5 X 1023 8x1014 Moderately Very fast fast at at all 20°C tempera­ tures

j

k

a Selected Values of Chemical Thermodynamic Properties, N.B.S. Technical Note 270-3 (1968). b A. G. Sharpe, Halogen Chemistry (ed. V. Gutmann), Vol. 1, p. 13, Academic Press (1967). c T. G. Kujumzelis, Physik, Z. 39 (1938) 665. d J. C. Evans and G. Y.-S. Lo, Inorg. Chem. 6 (1967) 1483. e J. C. Morris, / . Phys. Chem. 70 (1966) 3798. f C. H. Cheek and V. J. Linnenbom, / . Phys. Chem. 67 (1963) 1856. * O. Haimovich and A. Treinin, Nature, 207 (1965) 185. h Calculated from electrode potentials (ref. b). 1 M. W. Lister, Canad. J. Chem. 34 (1956) 465, 479. J P. Engel, A. Oplatka and B. Perlmutter-Hayman, / . Amer. Chem. Soc. 76 (1954) 2010. k C. H. Li and C. F. White, / . Amer. Chem. Soc. 65 (1943) 335; M. L. Josien and C. Courtial, Bull. Soc. chim. France (1949) 374; M. W. Lister and P. Rosenblum, Canad. J. Chem. 41 (1963) 3013; O. Haimovich and A. Treinin, / . Phys. Chem. 71 (1967) 1941.

point (-39-6°C) contains 47% HOC1 and the solid phase is then HOCl,2H2067i. Hypobromous acid, however, is stable only in dilute solution ( < 7 % HOBr); hypoiodous acid is unstable at all concentrations. Decomposition The manner and rate of decomposition of hypohalous acids or hypohalite ions in solution are much influenced by the concentration, pH and temperature of the solution, by added salts, and by photolysis. The two fundamental, competing styles of decomposition are and

2HOX -> 2H + + 2 X - + 0 2

(or 2 0 X " -> 2X" + 0 2 )

3HOX^3H++2X-+XO3-

(or 3 0 X - - > 2 Χ - + Χ 0 3 " )

OXYACIDS AND OXYSALTS OF THE HALOGENS

1405

The free acids react more readily than the anions, so that hypohalites are most stable in basic solution. Because of the commercial importance (in bleaching and sterilizing pro­ cesses) of the decomposition reactions of hypochlorites, there have been extensive studies of the effect of various additives as catalysts, promoters and activators; catalysis increases the rate of oxygen-evolution but not of disproportionation. For the hypohalite ions in basic solution disproportionation is the prominent method of decomposition. All the equilibrium constants are very favourable (Table 62), and the rate of the reaction increases in the sequence ClO" < BrO~ < IO~. The disproportionation of hypochlorite ion is bimolecular in OC1 ~, and involves the intermediate agency of the C102 ~ ion; the mechanism is slow

OC1-+OC1-

>C102-+C1-

(1)

fast

oci -+cio 2 - —► cio3 - + a -

(2)

Disproportionation of OBr~ and OI~ is autocatalysed by the halide ion formed; halite ion is again an intermediate, and solid bromites can be prepared by the controlled decomposi­ tion of hypobromite solutions (p. 1413). Halite ion arises directly from hypohalite [as in eqn. (1) above], and also in a halide-catalysed reaction XO- + X - + H + -> [X 2 OH]-

xo-

► 2X- +HXO2

slow

The bromite and iodite ions probably degrade via reaction with hypohalite [cf. eqn. (2)], though there is also some evidence for the alternative fast reaction

xo 2 -+xo 2 - ->xo 3 -+xoPhotolysis of alkaline hypochlorite solutions produces chloride, chlorate and oxygen, together with traces of chlorine dioxide (from photolysis of the chlorite intermediate)672. Evolution of oxygen from hypohalite solutions is also catalysed by cobalt, nickel and copper ions673. The hypohalous acids decompose more rapidly than the anions, and are more prone to yield oxygen, especially in the presence of light or catalytic amounts of metal ions; oxygenevolution represents the exclusive mode of decomposition of hypoiodous acid. Solutions of hypobromous and hypoiodous acids rapidly discolour through formation of the free halogens and trihalide ions in subsequent reactions of the disproportionation or decompo­ sition products. B r 0 3 - + 5 B r - + 6 H + ->3Br 2 + 3H 2 0 H O I + I - + H+ ->I2 + H20 X2+X->X3(X = BrorI)

The rate of disproportionation of hypochlorous acid is proportional to the square of the HOC1 concentration; the proposed mechanism has chlorous acid as an intermediate, slow

2HOC1 HOCl + C10 2 -

► 2H + + C\ ~ + C10 2 -

fast

>H++C1-+C103-

Oxygen-evolution from hypochlorous acid is unimolecular in HOC1. Like their solutions, solid hypohalites decompose more or less readily at room tempera­ ture either by disproportionation or by oxygen-evolution; the reactions are encouraged by 672 G. V. Buxton and R. J. Williams, Proc. Chem. Soc. (1962) 141. 673 M. W. Lister, Canad. J. Chem. 34 (1956) 479.

1406

CHLORINE, BROMINE, IODINE AND ASTATINE: A. J. DOWNS AND C. J. ADAMS

heating and by irradiation. The stability of solid hypochlorites is an important factor in determining their usefulness in the storage and transport of "active chlorine" for bleaching processes (see below). Reactions The hypohalites exchange oxygen rapidly with labelled water674. Hypochlorites oxidize many inorganic substrates (Scheme 8); kinetic studies reveal that either OCl~ or HOC1 is the oxidant, both species rarely being active in the same reaction. Most of the reactions involve oxygen atom-transfer to the substrate; isotopic studies have shown that hypochlorite oxidation of N 0 2 ~ involves quantitative transfer of oxygen from chlorine to nitrogen674, but with S0 3 2 ~ some sulphate is formed via C1S03~. S032+H0C1

-+CISO3-+OH-

CISO3 " + 20H " -> SO42 " + Cl" + H 2 0

While hypobromites and hypoiodites are thermodynamically weaker oxidizing agents, their reactions, which follow pathways similar to those of OC1 ~, are more rapid. Hypochlorites convert iodates to periodates, but are reluctant to oxidize chlorates675; hypoiodite rapidly brings about both reactions, but none of the hypohalites oxidizes Br0 3 ~ 676. While the halogens disproportionate in alkaline solution to halide and hypohalite (at least initially), in acid solution hypohalites combine with halides to regenerate the halogen: H O X + X - + H + - * X2+H2O

The equilibria are favourably disposed (Cl, 2-4 x 103; Br, 1 -4 x 108; I, 5 x 1012), and rapidly established in the case of bromine and iodine, the reactions occurring during the decompo­ sition of HOBr and HOI. The reaction of HOC1 with Cl ~ is slow, but the acid oxidizes Br to Br2 and I~ to I 2 . Hypohalite anions react with basic nitrogen by forming N-X bonds. Thus, hypochlorites convert ammonia and related compounds to chloramines; in dilute equimolar solutions NH 3 and hypochlorite generate the unstable chloramine NH2C1, NH 3 + OC1" -> NH2C1 + OH -

while with excess hypochlorous acid nitrogen trichloride is the product677. The familiar "chlorine" odour of water which has been sterilized with hypochlorite is due to chloramines produced from bacteria. However, hypobromites oxidize amines quantitatively to nitrogen, a facility exploited in the analysis of urea: 30Br- + 2 0 H - +CO(NH 2 ) 2 -> 3Br" + C 0 3 2 " + N 2 4 - 3 H 2 0

With primary amides the intermediate bromoamide loses C0 2 in rearranging to a primary amine. The reaction of hypohalites with cyanide involves the halogen cyanide as an inter­ mediate : C N - +OC1 +H2O -> C1CN+20HC1CN+20H-> NCO- +C1- + H 2 0 674 M . Anbar and H. Taube, / . Amer. Chem. Soc. 80 (1958) 1073. 675 M . W. Lister and P. Rosenblum, Canad. J. Chem. 39 (1961) 1645. 676 o . Haimovich and A. Treinin, / . Phys. Chem. 71 (1967) 1941. 677 D. G. Nicholson, Kirk-Othmer Encyclopedia of Chemical Technology, 2nd edn., Vol. 5, p. 2, Interscience, New York (1964).

1407

OXYACIDS A N D OXYSALTS OF THE HALOGENS

(alkali) OIΓ I03_ (acid) I

(alkali) OBr", BK)3" (acid) Br

C0 3 2 -,N 2 ,N0 3 -

ci2,cio3"

AsO/ sor SCHEME 8. Reactions of hypochlorous acid and hypochlorites with inorganic substrates.

Notes for Scheme 8. Kinetic studies of hypochlorite reactions: a M. W. Lister and P. Rosenblum, Canad. J. Chem. 39 (1961) 1645. b M. Anbar and H. Taube, / . Amer. Chem. Soc. 80 (1958) 1073. 0 M. W. Lister and P. Rosenblum, Canad. J. Chem. 41 (1963) 3013. a J. Halperin and H. Taube, / . Amer. Chem. Soc. 74 (1952) 380. H. Taube and H. Dodgen, /. Amer. Chem. Soc. 71 (1949) 3330. f F. Emmenegger and G. Gordon, Inorg. Chem. 6 (1967) 633. * L. Farkas, M. Lewin and R. Bloch, / . Amer. Chem. Soc. 71 (1949) 1988. h Y. T. Chia and R. E. Connick, /. Phys. Chem. 63 (1959) 1518. 1 M. W. Lister, Canad. J. Chem. 34 (1956) 489. i M. W. Lister, Canad. J. Chem. 34 (1956) 465, 479. k M. W. Lister and P. Rosenblum, Canad. J. Chem. 41 (1963) 2727. 1 R. E. Connick, / . Amer. Chem. Soc. 69 (1947) 1509. The hypohalous acids are used in organic chemistry as aromatic and aliphatic halogenating agents; HOBr and HOI are normally generated in situ. The ease of aromatic halogenation increases in the order hypochlorite < hypobromite < hypoiodite, and is encouraged by lead or silver salts. In acid solution the active agent may be either the protonated species H 2 OX + or X 2 0 (p. 1344)678; phenol is converted rapidly to 2,4,6-tribromo- or 2,4,6-tri-iodophenol in alkaline solution. Chlorination of aliphatic compounds with HOCl proceeds by a free radical mechanism57^. HOCl -*HO-+Cl· RH+C1· -*R+HC1 HC1+HOC1 ^ H 2 0+C1 2 R+CI2 ->RC1+C1· 678

3195.

E. Berliner, /. Chem. Educ. 43 (1966) 124; C. G. Swain and D. R. Crist, /. Amer. Chem. Soc. 94 (1972)

1408

CHLORINE, BROMINE, IODINE AND ASTATINE: A. J. DOWNS AND C. J. ADAMS

Hypohalites cleave methyl ketones, forming carboxylate anions and a haloform. RCOCH3+3OX- -> R C 0 2 - +CHX3+2OH-

The iodoform test has been used for many years to identify and enumerate CH3CO-groups in organic compounds, but direct titration against hypobromites using Bordeaux indicator has recently been used to determine methyl ketones679. Br BuOBr

o HOBr

[Pr'OBr]

NaOBr

RNH, PhCH(OH)CHBrCO,H

RC02Na+CHBr3

SCHEME 9. Reactions of hypobromous acid and hypobromites with organic substrates.

The acids generally add across ethylenic double bonds as though they were the ion-pair X + OH -, though the reaction of HOC1 with hindered olefins may follow a radical pathway. Hydrolysis of chlorohydrins (formed from olefins and HOC1) is a useful route to a-glycols. HOBr and especially HOI (i.e. Br2 or I 2 with HgO) have been used to oxidize alcohols via decomposition of the hypohalite ester. Bleaching and Other Uses Hypochlorites are widely used as relatively cheap agents for bleaching fabrics and wood pulp, although they are more likely to damage fabrics than more expensive bleaches (e.g. acid chlorite). Hypobromites are faster and stronger bleaching agents than hypochlorites680, and are conveniently generated in situ by adding small amounts of alkali bromide to hypochlorite solutions. The bleaching action of both systems depends on three main reac­ tions: (1) Disruptive oxidation of coloured molecules. (2) Addition of HOX across olefinic functions. (3) Halogenation of saturated compounds. The effects of various activators, promoters and catalysts on hypochlorite bleaching pro­ cesses have been investigated. 679 M. H. Hashmi and A. A. Ayaz, Anal. Chem. 36 (1964) 384. 680 M . Lewin, Bromine and its Compounds (ed. Z. E. Jolles), pp. 704-713, Benn, London (1966).

£ "Bleach liquor"; solution of Ca(OCl)2 and CaCl2 vo Calcium hypochlorite; normally dried Ca(OCl) 2 ,2H 2 0; water < 1 % "Bleaching powder"; Ca(OCl)2,CaCl2,Ca(OH)2,2H20 "Tropical bleaching powder*' Ί Ca(OCl)2,CaCl2, "Supertropical bleaching powder" >Ca(OH) 2 ,2H 2 0 Water sterilizing powder J + CaO Chlorinated trisodium phosphate dodecahydrate Lithium hypochlorite; normally diluted with sulphates

"Liquid bleach'*; sodium hypochlorite solution, pH ^ 11

Name and composition

TABLE 63. SOME

35% 34% 30% 25% 3-5% 40%

5% 10% 85gl-i 70%

Approximate "available chlorine" content

Detergent manufacture. Processes where calcium is undesirable; sanitation of hard water; some dairy applications.

[More stable than ordinary "bleaching powder", especially in hot | climates.

General bleaching and sanitation.

Domestic bleaching and sanitation. Small-scale commercial bleaching (e.g. in laundries). Pulp and paper bleaching. Swimming pool sanitation.

Special uses

PREPARATIONS USED IN BLEACHING, ETC.

1410

CHLORINE, BROMINE, IODINE AND ASTATINE: A. J. DOWNS AND C. J. ADAMS

The oxidizing powers of oxychlorine bleaches are indexed by the "available chlorine" content. This concept may be illustrated by considering the (possibly hypothetical) liberation of iodine from hydriodic acid. Comparison of the reactions C12+2HI

and

->2HC1 + I2

LiOCl + 2HI -► LiCl +12 + H20

shows that 1 mol of I 2 is liberated by 70-92 g of chlorine and also by 58-4 g of anhydrous lithium hypochlorite. The "available chlorine" content of LiOCl is defined as the weight of chlorine which liberates the same amount of I 2 as a given weight of the compound; expressed as a percentage, the "available chlorine" content of lithium hypochlorite is (70-92/58-4) x 100, or 121%. Some of the hypochlorite preparations used for domestic and industrial bleaching and sterilization operations are listed in Table 63. Basic sodium hypochlorite solutions (pH ca. 10-5) are sold for domestic use, but solid hypochlorites, being more stable and more concentrated in "available chlorine", are more convenient in terms of storage and transport. The most widely used materials are the hydrated and basic forms of calcium hypochlorite (including the traditional "bleaching powder") made by chlorination of slaked lime; "bleaching powder" has now been largely superseded by products which have either enhanced stability ("tropical" or "supertropical bleaching powder") or a higher "available chlorine" content. The stability of bleaching powders is adversely affected by carbonates, transition metal ions and organic impurities; calcium oxide increases the stability by taking up excess water. Esters of Hypohalous Acids Primary, secondary and tertiary alkyl hypohalites are produced by the reaction of an alcohol with a hypohalous acid (which may be generated in situ by hydrolysis of a halogen): R O H + H O X -> R O X + H 2 0

Primary and secondary derivatives are unstable, readily expelling hydrogen halide while forming respectively an aldehyde or a ketone, RR'CHOX -> R R ' C = O + HX TABLE 64. SPECTROSCOPIC INVESTIGATIONS OF METHYL HYPOCHLORITE

Microwave spectrum* Measured for i2CH 3 035Cl, 12CH3037C1, i3CH 3 035Cl, 12CH2D035C1, 12CD3035C1, 12CD3037C1 r a v (C-H) = 1·099±0·018 A; 0 av (H-C-H) = 109-4± 1-8° KO-C1) = l-674±0-019Ä; ö(C-O-Cl) = 112·8±2·1° K O - C ) = 1-389 ± 0 0 2 8 A Quadmpole coupling constants (CH 3 035Q): eZQq = - 8 4 3 4 MHz η = 0-408 Barrier to internal rotation — 3060± 150 cal mol" 1 b Infrared spectrum Recorded for vapour a b

J. S. Rigden and S. S. Butcher, / . Chem. Phys. 40 (1964) 2109. R. Fort, J. Favre and L. Denivelle, Bull. Soc. chim. France, (1955) 534.

OXYACIDS A N D OXYSALTS OF THE HALOGENS

1411

and are intermediates in the oxidation of alcohols by hypohalous acids (or by the halogen in the presence of Ag 2 0 or HgO). Methyl hypochlorite has been characterized spectroscopically in the vapour phase (Table 64). Tertiary-alkyl hypochlorites667 (and hypobromites681) are yellow liquids (and orange solids) at ambient temperatures, at least in the absence of light; illumination induces decomposition to a ketone and an alkyl halide: hv

RR'RXOX-> R R ' C = 0 + R ' X

The ease of fragmentation has been studied as a function of the alkyl substituents682. Tertiary-butyl hypochlorite and hypobromite are efficient halogenating agents for aliphatic and allylic C-H bonds; the radical mechanism involves the t-butoxy radical as a chaincarrier683 : hv Me3COX

->Me 3 CO

+X·

Me3CO · + RH -> Me3COH + R · R · + Me3COX -> R X + Me3CO

Inorganic chlorination by Me3COCl has also been reported684: HNF 2 + BuOCl -> BuOH + C1NF2

Heterolysis of the O-X bond is a possible alternative in some reactions. Thus, tertiaryalkyl hypohalites add across olefinic double bonds in a polar fashion, e.g. Q

+ ButOBr

MeQH

-

Q ™ * ♦ Bu'OH

and can act as sources of positive halogen for substitution into aromatic compounds: BuOCl is an excellent reagent for the oxidation of sulphides to sulphoxides without concomitant formation of sulphones685. BukKJl

R2S

R'OH

> [RzSClHBuO]-

fc -Bu OH

Δ

> R2S(OR')Cl

> R2S=0 -R'CI

Fluoroalkyl Hypochlorites These compounds are considerably more stable than the parent aliphatic esters; they are prepared either (1) by the elimination of HF between an alcohol and OF 6 8 6 , ROH+C1F — t ROC1+HF [R = (CF 3 ) 3 C, CH 3 (CF 3 ) 2 C, (CF 3 ) 2 CH or CF 3 CH 2 ]

or (2) by the addition at -20°C of C1F to a carbonyl compound in a reaction catalysed by CSF5S3.687, BF 3 or AsF 5 6 8 8 : RR'C= O+C1F -^ RR'FCOCl [R = F, R' = F or CF3; R = CF3, R' = CF3 or CF2C1] 681 C . Walling and A . Padwa, / . Org. Chem. 2 7 (1962) 2976. 682 c . Walling and A . Padwa, / . Amer. Chem. Soc. 85 (1963) 1593. 683 c . Walling and J. A . McGuinness, / . Amer. Chem. Soc. 91 (1969) 2053. 684 K . O. Christe, Iriorg. Chem. 8 (1969) 1539. 685 c . R. Johnson and J. J. Rigau, / . Amer. Chem. Soc. 91 (1969) 5398. 686 D . E. Young, L. R. Anderson, D . E . Gould and W . B . F o x , / . Amer. Chem. Soc. 9 2 (1970) 2313. 687 c . J. Schack and W. Maya, / . Amer. Chem. Soc. 91 (1969) 2902. 688 D . E . Young, L . R. Anderson and W . B . F o x , Inorg. Chem. 9 (1970) 2602.

1412

CHLORINE, BROMINE, IODINE AND ASTATINE: A. J. DOWNS AND C. J. ADAMS

The action of F4S = 0 on CIF in the presence of CsF generates the related compound F5SOC1583,689. CF3OCI is a by-product in the fluorination of OCF 2 with C1F3 adsorbed on y-alumina690, and may also be obtained from OCF 2 and C120583. Various spectroscopic techniques have been applied to characterize the compounds. Their thermal stabilities decrease in the order CF3OCl > C2F50C1 > (CF3)2CF0C1 > SF5OCI; CF3OCI decomposes only slowly at 150°C following a bipartite reaction path which involves the CF3O radical : CF30C1->CF30+C1 2CF 3 0 · -> CF3OOCF3 CF3O· - > O C F 2 + F

(91 %) (9%)

By contrast, photolysis of CF3OCl affords OCF 2 as the only detectable fragment691. The compounds combine readily with free-radical sources (e.g. N 2 F 4 ), and add to carbon monoxide and sulphur dioxide to form respectively chloroformates and chlorosulphates. hv

C F 3 O C l + i N 2 F 4 -> CF 3 ONF 2 +£Cl 2 O CF3OCI + CO

->CF3OC

CF3OCI+S0 2

-> CF 3 0S0 2 C1

Cl

Thiohypohalous Acids Under carefully controlled conditions bromine reacts with H2S in chloroform or dichloromethane692 : Br 2 + H 2 S - > H B r + H S B r

HSBr can be "fixed" with NH 3 in the form of an ammonium salt, which is stable at low temperature but decomposes rapidly in aqueous solution, NH 4 SBr -> S + N H 4 + + Br "

or on warming to room temperature. HSI has been similarly obtained, and SBr ~ is believed to be an intermediate in the oxidation by bromine of NaHS in ethereal solution.

(C) H A L O U S A C I D S A N D H A L I T E S

Introduction The acids HX0 2 and the anions X 0 2 ~ (X = Cl, Br or I) are intermediates in the disproportionation of HOX and OX ~ respectively, as well as in the oxidation of halides by halates. Chlorous acid, itself the least stable of the oxyacids of chlorine and known only in aqueous solution, is the best characterized of the halous acids; the existence of HBr0 2 and HI0 2 as other than transient species is still questionable. C102 ~ and Br0 2 " persist in neutral or alkaline solution if impurities are absent, and stable salts containing these ions may be isolated; I 0 2 ~ decomposes rapidly under all conditions. 689 c . J. Schack, R . D . Wilson, J. S. M u i r h e a d a n d S. N . C o h z , / . Amer. Chem. Soc. 9 1 (1969) 2907. 690 R . Veyre, M . Q u e n a u l t a n d C . E y r a u d , Compt. rend. 268C (1969) 1480. 691 K . O . Christe a n d D . P i l i p o v i c h , / . Amer. Chem. Soc. 9 3 (1971) 5 1 . 692 M . Schmidt a n d I . L ö w e , Angew. Chem. 7 2 (1960) 79.

OXYACIDS AND OXYSALTS OF THE HALOGENS

1413

Sodium chlorite is manufactured on a large scale, being extensively used as a bleaching agent (in acid solution) and as a source of chlorine dioxide for water purification and tallow bleaching. Both in the solid state and in solution the decomposition of chlorites produces chlorine dioxide. Under normal conditions chlorite solutions do not evolve C102 in dangerous amounts; however, explosive concentrations of C102 may result if acid is dropped or spilled on to solid chlorites, or if the salts are heated. Appropriate precautions should be taken. For reviews of the chemistry of the halogen(III) oxyacids and their salts the reader is referred to the following: chlorous acid and the chlorites572'574'575**»576»693'694; bromous acid and the bromites575b>576'668'695; iodous acid and the iodites575b>576»616. Preparation Halous Acids Chlorous acid is formed (together with HCIO3) in the decomposition of aqueous solutions of chlorine dioxide. Solutions of the acid were first obtained by reducing chloric acid with tartaric acid576, but the best method of preparation is to treat suspensions of barium chlorite with sulphuric acid: Ba(C102)2+H2S04 -> B a S 0 4 + 2 H C 1 0 2

Unstable solutions containing appreciable amounts of HBr0 2 695 or HI0 2 616 allegedly result from the reaction of aqueous silver salts with an excess of halogen; other acids (HX, HOX and HXO3) are also present in these solutions, among which the halous acids are apparently recognized by their characteristic oxidizing properties. Chlorites Solutions from which alkali or alkaline earth chlorites may be crystallized are best formed by reducing C10 2 : both industrially574 and in the laboratory, peroxides are con­ venient reducing agents (usually H 2 0 2 in conjunction with a solution or suspension of the metal hydroxide or carbonate). 2 C 1 0 2 + O 2 2 " -► 2 C 1 0 2 " + 0 2

Other reductants which have been used or suggested include organic matter, amines, nitrites, sulphur compounds, iodides and sodium amalgams. Some powdered metals of intermediate electropositive character (e.g. AI, Cd or Zn) unite directly with chlorine dioxide in aqueous solution (p. 1370); many more chlorites may be made metathetically from Ba(C102)2 and the appropriate metal sulphate. Bromites Aqueous solutions containing Br0 2 - can be obtained by the controlled disproportionation of cold concentrated alkaline hypobromite solutions 695-697 ; when the optimum bromite concentration is achieved, residual OBr ~ is destroyed by the addition of ammonia 693 C. C. Addison, Mellor's Comprehensive Treatise on Inorganic and Theoretical Chemistryt Supplement II, Part I, pp. 569-576, Longmans, London (1956). 694 A . S. Chernyshev, V. V. Shtutser and N . G. Semenova, Uspekhi Khim. 2 5 (1956) 91. 695 p . j . M. Radford, Bromine and its Compounds ( e d . Z . E . Jolles), pp. 159-162, Benn, London (1966). 696 H . Fuchs and R. Landsberg, Z. anorg. Chem. 372 (1970) 127. 69? M . Sediey, Fr. Addn. 92,474 and 92,491 (1968); Chem. Abs. 71 (1969) 126609q, 126610h.

1414

CHLORINE, BROMINE, IODINE AND ASTATINE! A. J. DOWNS AND C. J. ADAMS

or acetone. From such solutions bromites of lithium, sodium, potassium and barium may be crystallized. Anhydrous LiBr0 2 698 and Ba(Br02)2 6 " have been prepared by dry reactions: 190o 2LiBr0 3 + LiBr Ba(Br03>2

> 3LiBr0 2 250°

> Ba(Br0 2 ) 2 + 0 2

Properties Chlorous Acid Chlorous acid is a moderately strong acid (Table 65). In concentrations of only a few grams per litre it soon decomposes at room temperature, discolouring as chlorine dioxide is formed (see below). Acid solutions of chlorite are potentially stronger oxidizing agents than are alkaline solutions. TABLE 65. PROPERTIES OF AQUEOUS CHLOROUS ACID

Δ#/° for HC10 2 (undissoc.) AGf° for HC10 2 (undissoc.) 5° for HC10 2 (undissoc.) Dissociation constant, Ka Reduction potentials in acid solution0 £°(C10 3 " /HC102) £°(C10 2 /HC10 2 ) £°(HC10 2 /HOCl) £°(HC102/C1 ~ ) Ultraviolet-visible absorption spectrum

- 1 2 - 4 kcal m o l - i a + 1 -4 kcal mol" i a 45 0 cal deg ~ i mol" i 1 1 x 10 ~ 2 b +1-21 volts + 1 -27 volts + 1 -64 volts +1*57 volts d

a

Selected Values of Chemical Thermodynamic Properties, N.B.S. Technical Note 270-3 (1968). Consensus of the following individual studies: (i)Potentiometrictitration of C10 2 " withHCl; Ka = 1 0 7 x 1 0 - 2 ; G. F. Davidson, / . Chem. Soc. (1954) 1649. (ii) Spectrophotometric; Ka = 1-15x10-2; D. Leonesi and G. Piantoni, Ann. Chim. {Rome), 55 (1965) 668. (iii) pH titration of Ba(C102)2 with H 2 S0 4 ; Ka = 1-10x10-2; B . Barnett, Ph.D. Thesis, University of California (1935). (iv) Unspecified method; Ka = 101 x 10-2; M . W. Lister, Canad J. Chem. 30 (1952) 879. c A. G. Sharpe, Halogen Chemistry (ed. V. Gutmann), Vol. 1, p. 13, Academic Press (1967). d D. Leonesi and G. Piantoni, Ann. Chim. {Rome), 55 (1965) 668. b

Halites The chlorite and bromite ions have been studied spectroscopically in solution (Table 66) and in crystalline salts (Table 67); AgOC>2 and NH 4 C10 2 have been investigated crystallographically. The O-Cl-O angle (111°) is close to the expected tetrahedral value (cf. C103~, 110°andClO 2 , 118°). There are large and unexplained discrepancies between the vibrational frequencies attributed to Br0 2 ~ in the infrared spectra of the solids LiBr0 2 and Ba(Br02)2 and those observed in the Raman and infrared spectra of Br0 2 " in aqueous solution and in NaBr0 2 ,3H 2 0; the latter study is more plausible, since the frequencies observed for the lithium and barium salts are high in comparison with those of C10 2 ", and yield on analysis an unusually large valence force constant for Br-O stretching (4-75 mdyne A" 1 ) 6 ". Decomposition of Chlorites Alkaline solutions of sodium chlorite remain unchanged for up to one year if light is 698 p . Hagenmuller and B. Tanguy, Compt. rend. 260 (1965) 3974. 699 B . Tanguy, B. Frit, G. Turrell and P. Hagenmuller, Compt. rend. 264C (1967) 301.

OXYACIDS AND OXYSALTS OF THE HALOGENS

1415

TABLE 66. PROPERTIES OF AQUEOUS HALITE IONS

Colour Thermodynamic properties at 298°K A#/°(kcalmol-i) AG/°(kcalmol-i) S^caldeg-imol"1) Reduction potentials in alkaline solution* E°(X03-IX02) (volts) Ε ° ( Χ 0 2 / Χ 0 2 - ) (volts) £ ° ( X 0 2 - / X O - ) (volts) £ ° ( X 0 2 - / X - ) (volts) Raman spectrum Fundamental frequencies (cm - 1 ) *Ί(«Ι) v2{fli)

"3(W Force constants (mdyne Ä" 1 ) fr frr

fel*

Mr

Infrared spectrum Ultraviolet-visible absorption spectrum ^max(nm)

cio 2 -

Br02-

Colourless

Yellow

-15-9a +4-la 24-2 a +0-33 + 116 +0-66 +0-77 c 790 400 [840]e f 4-26 011 0-52 002 g 260

d 709 324 680 d 4-2

d h 290

a Selected Values of Chemical Thermodynamic Properties; N.B.S. Technical Note 270-3 (1968). b A. G. Sharpe, Halogen Chemistry (ed. V. Gutmann), Vol. 1, p. 13, Academic Press (1967). c J. P. Mathieu, Compt. rend. 234 (1952) 2272. d J. C. Evans and G. Y.-S. Lo, Inorg. Chem. 6 (1967) 1483. e Estimated from spectra of solid chlorites. f H. Siebert, Anwendungen der Schwingungsspektroskopie in der Anorganischen Chemie, p. 50, Springer-Verlag (1966). β W. Buser and H. Hänisch, Helv. Chim. Ada, 35 (1952) 2547; H . L. Friedman, / . Chem. Phys. 21 (1953) 319; D. Leonesi and G. Piantoni, Ann. Chim. {Rome), 55 (1965) 668. h O. Amichai and A. Treinin, / . Phys. Chem. 74 (1970) 830.

excluded; even with boiling no decomposition occurs. The stability decreases as the pH falls, cold neutral solutions being stable in the dark, but degrading slowly if heated. Acid solutions decompose at measurable rates, especially if the pH is less than 4. The decomposition of chlorous acid may be expressed by three equations, none of which on its own accurately describes the reaction: (i)

3HC102->3H++2C103-+C1-

(ii)

5HC10 2 -► 4C10 2 +C1" + H

(iii)

HC10 2 -> Cl" + H + + 0 2

+

+2H20

AG° = - 3 3 - 2 kcal mol"* AG° = - 3 4 - 5 kcal mol"i AG° = - 29-5 kcal mol -1

Oxygen-evolution (iii) is a minor process, accounting for less than 3% of the chlorite consumed. The simple disproportionation to chlorate and chloride (i) is not observed, and in 2 M perchloric acid the stoichiometry of the disproportionation of HCIO2 initially

1416

CHLORINE, BROMINE, IODINE AND ASTATINE! A. J. DOWNS AND C. J. ADAMS TABLE 67. PROPERTIES OF SOLID HALITES

X-ray diffraction NH4C102& (at - 35°C)r(Cl-0) = 1·57±0·03Α; 0O_C.t_0 = 110-5 + 1-4° AgCl02h r(Cl-O) = 1-55 ± 0 0 5 Ä; 0O-ci-o = 111 ±2\° Vibrational spectra "3(61) vi(a\) ^l(«l) (cm-i) (cm-i) (cm-i) c 844 Raman 402 786 NaC10 2 ,3H 2 0 860 Infrared 1 —330 720 680 NaBr0 2 ,3H 2 0 Raman d 800 400 775 LiBr0 2 Infrared e 800 400 775 Infrared e Ba(Br0 2 ) 2 3 s 5C1 quadrupole resonance spectra Resonance frequencies: AgC10 2 , 5408 MHz; NaC10 2 , 51 -82 MHz. Fluorescence1 Emission11 X-ray spectra: NaC10 2 E.S.C.A.1 a

R. B. Gillespie, R. A. Sparks and K. N. Trueblood, Acta Cryst. 12 (1959) 867. J. Cooper and R. E. Marsh, Acta Cryst. 14 (1961) 202. J. P. Mathieu, Compt. rend. 234 (1952) 2272. d J. C. Evans and G. Y.-S. Lo, Inorg. Chem. 6 (1967) 1483. e B. Tanguy, B. Frit, G. Turrell and P. Hagenmuller, Compt. rend. 264C (1967) 301. f Includes data for AgC10 2 and Pb(C10 2 ) 2 ; C. Duval, J. Lecomte and J. Morandat, Bull. Soc. chim. France (1951) 745. « J. L. Ragle, / . Chem. Phys. 32 (1960) 403. h V. I. Nefedov, / . Struct. Chem. 8 (1967) 919. 1 A. Fahiman, R. Carlsson and K. Siegbahn, Arkiv. Kemi, 25 (1966) 301. 3 J. A. Bearden, Rev. Mod. Phys. 39 (1967) 86; D. S. Urch, / . Chem. Soc. (A) (1969) 3026. b

c

approximates to equation (iv)700: (iv) 4HC102 -> 2C102 + CIO3 + Cl" + 2H + + H20 However, reaction (ii) is catalysed by C\ ~, so that the rate of formation of CIO3 ~ relative to CIO2 decreases as the reaction proceeds; if appreciable amounts of chloride are present, reaction (ii) predominates. A very small amount of CIO2 also results from the interaction of chlorous acid with chlorate701, viz. HC102 + H + + CIO3 - -+ 2C102 + H20 The overall rate law has the form

~"[HC1°2]

wao2v+kuHao>m]2

dt

K+[C\~]

and the following mechanisms have been proposed. (a) Uncatalysed reaction HC10 2 +HC10 2

-> HOC1+ H + + CIO3 " (rate-determining)

HOC1 + HC10 2

->

HOCl+H++Cr Cl 2 + HC10 2 C1-C1(^ 2 Cl-Cl(^ 700 701

[-<]■ Cl-CT

+ H20

^C12+H20 -*

C\-C\/

+ H + + Cl "

+H20->C1-+C103-+2H

+

->C1 2 +2C10 2

R. G. Kieffer and G. Gordon, Inorg. Chem. 7 (1968) 235, 239. H. Taube and H. Dodgen, / . Amer. Chem. Soc. 71 (1949) 3330.

OXYACIDS A N D OXYSALTS OF THE HALOGENS

1417

(b) Chloride ion catalysed reaction HCIO2+CI^[HC1202-] [HCI2O2 " ] + Cl" -> products (rate-determining) At low hydrogen ion concentrations the following path may also assume importance701 : HC102+C102- -> [HCIO2-CIO2-3 -> products Chlorous acid solutions can be stabilized by the addition of H 2 0 2 , which regenerates C102 ~* from chlorine dioxide. Pyrolysis of sodium chlorite proceeds with disproportionation to chlorate and chloride702: 260°C

3NaC10 2

► 2NaC10 3 + NaCl

While disproportionation is also the process initially observed on heating chlorites of heavier metals, many of these salts decompose explosively at higher temperatures or on rapid heating, probably via the evolution of C102 703. Decomposition of Bromites and Iodites The mechanisms by which aqueous Br0 2 ~ and I 0 2 ~ ions decompose have been much less fully studied than those of the chlorite reactions; though the primary processes are presumably analogous to those in the disproportionation of chlorite, there is some evidence for the occurrence in alkaline media of the simple reaction676»704 X 0 2 - + X 0 2 - ~> X O 3 - + X O -

(X = Br or I)

There are no reports that heating solid bromites occasions disproportionation. LiBr02 (at 225°C)7°5 and Ba(Br02)2 (at 220°C)699 lose oxygen in forming the metal bromide, whereas after dehydration Mg(Br02)2,6H20 decomposes to magnesium oxide706. Crystalline bromites are stable at room temperature, even if illuminated. Reactions Fluorination of sodium chlorite produces low yields of CIOF3 641. Chlorine and hypochlorous acid oxidize chlorites to C102 or CIO3 _ or to a mixture of both. The reaction of chlorine with solid sodium chlorite comprises an expedient synthesis of C102. 2 N a C 1 0 2 + C l 2 -> 2NaCl+2C10 2

In acidic solution chlorine reacts more rapidly and produces less chlorate than does hypochlorous acid, although the product distribution also depends on the ratio of the reactants; excess chlorite encourages the formation of C102 707. The reaction HOC1 + HCIO2 -> CIO3- + C 1 - + 2 H +

is the second stage in the disproportionation of hypochlorites (p. 1405), and similar reactions have been postulated between the other halites and hypohalous acids. Chlorous acid is oxidized by permanganate to chloric acid. 702 F . Solymosi, T . Bansagi a n d E . Berenyi, Magy. Kern. Foly. 7 4 (1968) 2 3 ; Chem. Abs. 6 8 (1968) 51465s. 703 F . Solymosi, Z. physik. Chem., N.F. 5 7 (1968) 1. 704 c . H . C h e e k a n d V . J . L i n n e n b o m , / . Phys. Chem. 6 7 (1963) 1856. 70s R . D i a m e n t , Compt. rend. 2 6 4 C (1967) 589. 706 R . D i a m e n t a n d M . Sediey, Compt. rend. 2 6 8 C (1969) 1243. 707 F . E m m e n e g g e r a n d G . G o r d o n , Inorg. Chem. 6 (1967) 6 3 3 .

1418

CHLORINE, BROMINE, IODINE AND ASTATINE: A. J. DOWNS AND C. J. ADAMS

In acid solution chlorites and bromites oxidize iodide ion to iodine, which is further oxidized to IO3 ~. Bromites rapidly oxidize cold alkaline arsenite solutions, with which chlorites react only slowly. The oxidation of transition metal ions by chlorite has also been explored708. Uses Sodium chlorite is widely used in acid solution as a bleaching agent for textiles574; typically pH values of 3-5 have been used. The nature of the bleaching reactions is uncertain; CI2O3 is a possible active intermediate, although oxidation of coloured impurities by oxygen and chlorine dioxide (decomposition products of HC102) is also important. Recent improvements in the efficiency of chlorite bleaching have involved the use of solutions of relatively high pH (to hinder the escape of C102) and of various activators. Alkaline bromite solutions are used in desizing cellulose fabrics; this treatment removes starches which are added to fabrics to facilitate weaving but which interfere with the penetration of reagents used in subsequent bleaching and dyeing processes.

(D) H A L I C A C I D S A N D H A L A T E S

Introduction

In 1658 J. R. Glauber claimed that he could convert hydrochloric acid into "nitric acid", and common salt into "saltpetre"; the substances he produced were almost certainly chloric acid and sodium chlorate576. Towards the end of the eighteenth century several workers investigated crystalline solids generated by the interaction of the recently discovered chlorine with alkaline materials; the chemical constitution of these salts (chlorates) was recognized by Berthollet. In 1814 Gay Lussac obtained solutions of chloric acid by treating barium chlorate with sulphuric acid. The oxidation of iodine, suspended in water, to iodic acid was reported in 1813 by Davy (who used C102) and by Gay Lussac (who used chlorine); the latter author also investigated the disproportionation of iodine (forming iodide and iodate) in alkaline media. Barium bromate and bromic acid were described by Balard in 1826. Calcium iodate occurs naturally in the form of the minerals lautarite and dietzeite, found, for example, in the Caliche deposits of Chile (see p. 1127). The halates are strong oxidizing agents in acid solution, and in so acting are usually reduced to the corresponding halide ions. As judged by the redox potentials (Table 69), their oxidizing powers decrease in the order bromate ^ chlorate > iodate, but the rates of reaction follow the sequence iodate > bromate > chlorate. The couples (Br0 3 - /Br -) and (Br0 3 "/ jBr 2 ) have approximately the same potentials in acid solution as the (Mn0 4 ~/Mn2 +) couple (+1*51 volts). The solid compounds form spontaneously inflammable or explosive mixtures with oxidizable material, and acidified chlorates are liable to release C102 in dangerously high local concentrations. Chlorates are important precursors in the manu­ facture of chlorine dioxide and perchlorates, and are used in making matches and fireworks; bromates and iodates have only minor commercial applications, but are useful analytical reagents. 708 G . Gordon and F. Feldman, Inorg. Chem. 3 (1964) 1728; B. Z. Shakhashiri and G. Gordon,/. Amer. Chem.Soc. 91 (1969) 1103.

1419

OXYACIDS AND OXYSALTS OF THE HALOGENS 572 5 5

576

There have been several reviews of chloric acid and the chlorates » ? *' »™^^ bromic acid and the bromates575t>,576,7ii,7i2 a n d iodic acid and the iodates575b,576,624,7i3. Formation and Preparation The best routes to many chlorates, bromates and iodates (not excluding the acids) involve metathetical or neutralization reactions in aqueous solution. Lists of individual reactions will be found elsewhere572»576»709'711»713, but a general method involves the action of a sulphate on the appropriate barium halate: Ba(X03)2+M2S04 -> 2MX03+BaS04 The anions themselves are usually produced either (i) by disproportionation of the parent halogen in alkaline solution or (ii) by chemical or electrolytic oxidation of the parent halogen or of the corresponding halide ion. Only the most important processes are described below. Halic Acids Treating Ba(C103)2 or Ba(Br0 3 ) 2 with the stoichiometric amount of sulphuric acid remains the best means for obtaining solutions of HC10 3 or HBr0 3 . Ba(X03)2+H2S04 -* BaS04+2HX03 Although amenable to the same mode of preparation, iodic acid is more readily synthesized by the oxidation of an aqueous suspension of iodine, either electrolytically or chemically (with fuming nitric acid); crystallization of moderately acid solutions deposits α-ΗΙ0 3 , but HI3Og emerges from solution in concentrated ( > 5 M) acid. Laboratory Synthesis of Halates The disproportionation of the halogens in basic solution offers a facile route to halate ions: 3X2+6OH- - * X 0 3 - + 5 X - + 3 H 2 0

Disproportionation is speeded at higher temperatures, so that the free halogen is added to a warm (50-80°C) solution of a hydroxide or a carbonate; in the case of bromates, boiling of the solution is recommended to hasten the conversion of intermediate hypobromites and bromites to bromate. Fractional crystallization readily separates the less soluble halate from the accompanying halide. It is also possible to chlorinate suspensions of metal oxides or hydroxides. To synthesize bromates and iodates, oxidation of the halogen or halide is often more economical than disproportionation of the halogen. Thus bromides are transformed into bromates by hypochlorites in aqueous solution (conveniently executed by passing chlorine into an alkaline bromide solution), or by molten chlorates. Aqueous chlorate oxidizes bromine only slowly. 709 c . C. Addison, Mellows Comprehensive Treatise on Inorganic and Theoretical Chemistry, Supplement II, Part I, pp. 576-596, Longmans, London (1956). 710 T. W. Clapper and W. A. Gale, Kirk-Othmer Encyclopedia of Chemical Technology, 2nd edn., Vol. 5, pp. 50-61, Interscience, New York (1964). 711 B. Cox, Mellows Comprehensive Treatise on Inorganic and Theoretical Chemistry, Supplement II, Part I, pp. 753-773, Longmans, London (1956). 712 p. J. M. Radford, Bromine and its Compounds (ed. Z. E. Jolles), pp. 162-178, Benn, London (1966). 713 G. J. Hills, Mellor's Comprehensive Treatise on Inorganic and Theoretical Chemistry, Supplement II, Part I, pp. 874-895, Longmans, London (1956).

1420

CHLORINE, BROMINE, IODINE AND ASTATINE: A. J. DOWNS AND C. J. ADAMS

Iodates may be prepared directly by heating alkali-metal iodides to ca. 600°C under high pressures of oxygen. Chlorates (or bromates) oxidize iodine to the corresponding iodates, e.g. I 2 +2NaC10 3 -> 2NaI0 3 + Cl2

Hypochlorite may not be used since this causes further oxidation of iodine(V) to iodine(VII)714. Iodates also result from the thermal decomposition of periodates (p. 1457). Manufacture of Chlorates and Bromates Chemical methods for manufacturing chlorates and bromates involve the halogenation of an alkali (as described above). For chlorate this process has largely been superseded by electrolytic oxidation of chloride, but some bromate is still produced by dissolving bromine in a warm alkaline solution or suspension, since the bromide simultaneously formed is a desirable by-product. Brine is the common electrolyte in cells for chlorate production710»715. No diaphragm separates the electrodes (a steel cathode and, usually, a graphite anode) which are situated close together. Under optimum conditions present cells have a current efficiency of 80-90% for the overall cell process Cl - + 3H 2 0 ■-> C10 3 " + 3H 2

The primary electrode reactions are oxidation of Cl ~ to CI2 at the anode: Cl" ->£Cl2+e and production of hydroxide ions at the cathode: H 2 0+e->£H 2 +OHDiifusion in the cell allows the chlorine to disproportionate: Cl2+20H - -* Cl -+ OCl -+ H20 Chlorate then arises either by disproportionation of hypochlorite (p. 1405) or by direct oxidation of OCl ~ at the anode. Losses occur through extraneous oxidation and reduction processes at the electrodes, by decomposition of hypochlorite (to chloride and oxygen), and by consumption of the anode. The electrochemical production of bromates follows similar lines712. However, the electrolyte is now a solution of bromine in alkali, so that some OBr - is initially present, and anodes coated with PbC>2 are employed; reduction of OBr" at the cathode is prevented by the addition of small amounts of dichromate. Properties and Structural Chemistry Halic Acids HCIO3 and HBr0 3 can be obtained only in aqueous solution. Dilute solutions decom­ pose on boiling, but they may be concentrated to syrupy liquids under vacuum at room temperature. The limiting concentrations for stability are: HCIO3 ca. 40%, HBr0 3 ca. 50%; decomposition paths are discussed below. Iodic acid solutions crystallize on evaporation, and heating solid HIO3 occasions only dehydration, producing ultimately its anhydride, I2O5. 714 H. H. Willard, Inorganic Syntheses, Vol. I (ed. H. S. Booth), p. 168, McGraw-Hill (1939). 715 J. C. Schumacher, / . Electrochem. Soc. 116 (1969) 68C.

OXYACIDS AND OXYSALTS OF THE HALOGENS

1421

Both chloric and bromic acids are strong acids (pKa £ 0); by contrast, iodic acid is only moderately strong (pKa, 0-8), and the presence of the undissociated molecule may be detected in solution by spectroscopic methods (Table 68).

2

4

M[CrI06](M=K, Rb,CsorNH4)

SCHEME 10. Reactions of some iodates. TABLE 68. PROPERTIES OF IODIC ACID AND ANHYDRIODIC ACID HI3Oe

HIO3 Melting point (°C)a Dehydration temperature (°C)a Thermodynamic functions at 298°Kb A//>°(c)(kcalmol-i) A//>°(undissoc,aq) (kcal mol~ *) AG/°(undissoc,aq) (kcal mol"1) S°(undissoc,aq) (cal deg"1 mol"1) Acidity Dissociation constant, Ka Hammett acidity function, Ho Crystal structure X-ray diffraction Neutron diffraction

Dimensions ofHOlOi molecule KI=0)(A) r(I-OH)(A) 0O-I-OH

0o»i»o

Spectroscopic investigations Raman spectrum Infrared spectrum Mlquadrupole resonance spectrum e2Qq(MHz) 1

Mass spectrum

HI 3 o 8

ca.110 ca. 100 (->HI 3 0 8 ) 1 ί

CÖ.200(->I2O5)

-92-2 d

-55-4° -50-5* -31-7 d 39-9* 0-157* exptf calc* a-HIOa11 a-HKV Contains HOIO2 molecules

1

Contains HOIO2 and I2O5

molecules

h

t

1-80,1-82 1-89 95-7°, 98-2° 101-4°

1-79,1-81 1-90 94-2°, 95-4°

solid*·«1 aq soln

solidk

990

solid*·*

s o l i d »,k,m,n

solid0 1126-9 0-4505 p vapour

1

1422

CHLORINE, BROMINE, IODINE AND ASTATINE: A. J. DOWNS AND C. J. ADAMS

Table 68 {cont.) » K. Seite and A. Kjekshus, Acta Chem. Scand. 22 (1968) 3309. b For redox potentials see Table 69. c Estimated from the heat of reaction of HIO3 with aqueous N2H4; P. B. Howard and H. A. Skinner, / . Chem. Soc. (A) (1967) 269. d Selected Values of Chemical Thermodynamic Properties, p. 40, N.B.S. Technical Note 270-3 (1968). e Conductimetric, potentiometric and kinetic-salt effect investigations of solutions < 0 1 M ; A. D . Pethybridge and J. E. Prue, Trans. Faraday Soc. 63 (1967) 2019; includes critical survey of earlier measurements of Ka for HIO3. f Measured for solutions < 6 M using o- and /Miitroaniline as indicators; J. G. Dawber, / . Chem. Soc. (1965) 4111. « J. G. Dawber, / . Chem. Soc. (A) (1968) 1532. h Error limits for distances ± 0 04 A, for angles ± 5 ° ; M. T. Rogers and L. Helmholz, / . Amer. Chem. Soc. 63 (1941) 278. 1 Error limits for distances ±0-01 Ä, for angles ±0-6°; Y. D . Feikema and A. Vos, Acta Cryst. 20 (1966) 769. J r ( I — O H ) = l - 8 9 9 ± 0 0 1 1 Ä , r ( I = O ) = 1·780±0·010Α, 1·816±0·010Α; B. S. Garrett, U.S. Atomic Energy Comm. ORNL-1745 (1954); Chem. Abs. 49 (1955) 5064b. k P. M. A. Sherwood and J. J. Turner, Spectrochim. Acta, 26A (1970) 1975. 1 G. C. Hood, A. C. Jones and C. A. Reilly, / . Phys. Chem. 63 (1959) 101. m Includes data for HIO3 and DIO3; W. E. Dasent and T. C. Waddington, / . Chem. Soc. (1960) 2429. n L. Couture, M. Krauzman and J. P. Mathieu, Compt. rend. 269B (1969) 1278. 0 Single crystal study, including Zeeman effect; R. Livingston and H. Zeldes, / . Chem. Phys. 26 (1957) 351. p M. H. Studier and J. L. Huston, / . Phys. Chem. 71 (1967) 457. q Includes data for HIO3 and DIO3; J. R. Durig, O. D. Bonner and W. H. Breazeale, / . Phys. Chem. 69 (1965) 3886.

Investigations of the system I2O5-H2O have shown two modifications of iodic acid and the presence of an intermediate phase, anhydriodic acid Ηΐ3θ8 618. Both a-HI0 3 and J8-HIO3 are orthorhombic, differing only in the ratio of the crystallographic axes. Singlecrystal diffractometric studies (summarized in Table 68) reveal HOI0 2 molecules in 0C-HIO3, and both HOI0 2 and J 2 0 5 molecules in HI 3 0 8 (which is therefore to be formulated as HIO3J2O5); in both compounds there is evidence of hydrogen-bonding and of significant I· · O intermolecular contacts (at 2-5-3Ό Ä) which result in distorted octahedral coordina­ tion about the unique iodine atom in 0C-HIO3 and about two of the iodine atoms in HI3O8. The third iodine atom in HI 3 0 8 has 4 intermolecular I · · · O contacts, giving irregular hepta-coordination. Polymerization of Iodic Acid The physical (conductimetric and potentiometric) properties of aqueous iodic acid solutions are consistent with the formation in moderately concentrated solution of the anion [H(I03)2] ~ 716, having a stability constant of 4 1 mol _1 as defined by the equation HIO3+IO3- ^ [ H ( I 0 3 ) 2 1 -

The behaviour of the acidity function of iodic acid in concentrated solution indicates further association to furnish [I03(HI03)n] - (n > 1)717, Hydrogen biiodate and dihydrogen triiodate salts, M I H(I0 3 ) 2 and MIH2(I03)3, may be crystallized from solutions containing ΜΊΟ^ and excess HIO3 718. KH(I0 3 ) 2 , which is used as an alkalimetric standard, suffers dehydration at 105°C; infrared studies of its 716 A . D . Pethybridge a n d J . E . P r u e , Trans. Faraday Soc. 6 3 (1967) 2019. 717 j . G . D a w b e r , / . Chem. Soc. (1965) 4 1 1 1 ; ibid. (A) (1968) 1532. 718 S. B . Smith, / . Amer. Chem. Soc. 69 (1947) 2285.

OXYACIDS AND OXYSALTS OF THE HALOGENS

1423

719

structure have proved inconclusive , while preliminary crystallographic data have been reported for the a- and y-modifications720. Heating these hydrogen iodates produces salts such as Κ2Η2ΐ6θι7, which contains I-O-I bonds721. Chlorates, Bromates and Iodates Alkali-metal chlorates and bromates crystallize with distorted NaCl-type lattices; however, the corresponding iodates adopt perovskite structures and display piezoelectric properties. The solubility of these compounds in water decreases in the order chlorate > bromate > iodate. The iodates of some tetravalent metal ions (e.g. Ce, Zr, Hf and Th) are sufficiently insoluble to afford a useful means of separation. Double salts between alkali-metal iodates and halides have also been investigated [e.g. K3C1(I03)2]. Iodates form complex polyanions with the oxyanions of other metallic and non-metallic elements, e.g. sulphates, selenates, tellurates, molybdates, tungstates, chromates and vanadates, and also iodato-complexes with some metal ions, e.g. [M(I03)Ö] 2 " (M = Mn, Ti and Pb). As expected with the presence of a non-bonding electron pair on the halogen atom, the XO3 ~ anions defined by diffractometric techniques in the salts listed in Table 70 are pyra­ midal, and usually have C?>v symmetry within the experimental error. Owing to the difficulty encountered in fixing oxygen atom positions, many early studies of bromates and iodates yielded only imprecise parameters which have not been included in Table 70. The CIO3 - and Br0 3 ~ ions are seen to be rather flat pyramids with O-X-O angles of ca. 106°; however, the iodate ion has smaller bond angles {ca. 97°), possibly indicating that there is less ^-character in the I-O bonds. As might be anticipated, the X-O bond lengths increase from chlorate to bromate to iodate. The HOI0 2 molecules in α-iodic acid and HI3O8 have two I-O bonds close in length to those of the iodate ion (ca. 1-80 A), but the I-OH bonds are longer (1-90 A); the bond angles are all close to 100° (Table 68). Significant intermolecular I· · Ό interactions are found in crystalline iodates and iodic acid, the distances (2-5-3-3 A) being appreciably less than the sum of the van der Waals' radii of iodine and oxygen (3-55 A); such interactions give rise to a trigonally distorted octahedral environment about the iodine in 0C-HIO3, LÜO3, NH4IO3 and Ce(I03)4,H20 and about two of the iodine atoms in H^Og, to hepta-coordination in Sr(I03)2,H20 and (about one of the iodine atoms) in HI3O8, and a square-antiprism of eight oxygens about the iodine in NaI0 3 and Ce(I03)4. The high viscosity of aqueous iodate solutions and the low mobility found for the IO3" ion imply that the ion is strongly solvated in solution722. Infrared and Raman spectroscopic results confirm that the XO3 - ions have in solution essentially the same pyramidal configuration as in solid derivatives; the spectroscopic features of the XO3 ~ anion in simple halates are somewhat modified by the crystal environ­ ment of the anion. However, the spectra of some bromates and iodates [including M(I03)2 (M = Pb or Hg), K 2 M(I0 3 ) 6 (M = Pb, Ti or Mn), Fe(I03)3, [Co(NH3)3(H20)(I03)2]I03 and M(Br03)3,4H20 (M = La, Pr or Nd)] display bands attributable to stretching motions of M - 0 - X 0 2 units rather than to the four fundamentals characteristic of the isolated XO3 anion. The salt NH 4 CrI0 6 contains the anion [03CrOI02] -, composed of a Cr0 4 tetra­ hedron and an IO3 pyramid sharing one corner; its dimensions are: r(Cr-Oterm) = 1-64 A; KCr-Obridge) = 1-90 A; r(I-Oterm) = 1-83 A; r(I-Obridire) = 2-06 A; Zl-O-Cr = 118° (see Table 70, ref. m). 719 W. E. Dasent and T. C. Waddington, / . Chem. Soc. (1960) 2429. 720 G. Argay, I. Naray-Szabo and E . Peter, / . Therm. Anal. 1 (1969) 413. 721 T. Dupuis, Mikrochim. Acta, (1962) 289. 722 H . A . Bent, Chem. Rev. 68 (1968) 615.

C.I.C. VOL II—XX

1424

CHLORINE, BROMINE, IODINE AND ASTATINE: A. J. DOWNS AND C. J. ADAMS

TABLE 69. PROPERTIES OF HALATE IONS IN AQUEOUS SOLUTION

Thermodynamic functions at 298°K AÄ>°(kcalmol-i) Ä AG/°(kcalmori) a ^(caldeg-imol"1)» Reduction potentials Acid solution £°(X04-7X0 3 -)(volts) a £°(H 5 X0 6 /X0 3 -)(volts)t> £°(X03-7X0 2 )(volts) b £ 0 (X0 3 -/iX 2 )(volts) a ' b E°(X0 3 -/X-)(volts)». b Alkaline solution E°(X04 " / X 0 3 * )(volts) a ' b £ 0 (X0 3 -/iX 2 )(volts)*. b E°(X03-IX-)(vo\ts)*>» Spectra, etc. Raman spectrum Infrared spectrum Fundamental frequencies (cm ~*) nifli) v2(ai) v*(e) v4(e) Force constants1 /rCmdyneA* 1 ) / ö /r2(mdyneÄ-i) Ultraviolet-visible spectrum* Amax(nm) Nmr spectrum 17 0 chemical shift (ppm with respect to H 2 0 ) 35 C1 chemical shift (ppm with respect to Cl~) Ionic mobility at 25°C, AQ (cm2 o h m - 1 g-ion" 1 )

cio 3 -

Br03"

-24-87 -1-90 38-8

-15-95 +4-55 38-6

+ 1-23

+ 1-76

io 3 -52-51 -30-20 28-3

ca. +1-7 + 1-15 + 1-46 + 1-44

+ 1-51 + 1-44

+ 1-20 + 109

+0-40 +0-47 +0-62

+0-93 +0-52 +0-61

+0-81 +0-20 +0-29

Cd.e.i M

d,f,g,i 1

d.e.h.i

i

I

i i

933 608 977 477

805 418 805 358

805 358 775 320

5-715 1025

5-068 0-633

5-16 0-49

<200

195

195

-287k

-297k

-10501 64-58 m

55-8 m

40-54*

a G. K. Johnson, P. N . Smith, E. H. Appelman and W. N. Hubbard, Inorg. Chem. 9(1970)119. b A. G. Sharpe, Halogen Chemistry (ed. V. Gutmann), Vol. 1, p. 13, Academic Press (1967). c C. Rocchiccioli, Ann. Chim. Series 13, 5 (1960) 999. d G. W. Chantry and R. A. Plane, / . Chem. Phys. 34 (1961) 1268. e S. T. Shen, Y. T. Yao and T. Wu, Phys. Rev. 51 (1937) 235. f M. Rolla, Gazz. Chim. Ital. 69 (1939) 779. * C. Rocchiccioli, Compt. rend. 249 (1959) 236. 11 G. C. Hood, A. C. Jones and C. A. Reilly, / . Phys. Chem. 63 (1959) 101. 1 D. J. Gardiner, R. B. Girling and R. E. Hester, / . Mol. Structure, 13 (1972) 105. 1 A. Treinin and M. Yaacobi, / . Phys. Chem. 68 (1964) 2487. k B. N . Figgis, R. G Kidd and R. S. Nyholm, Proc. Roy. Soc. 269A (1962) 469. 1 Y. Saito, Canad. J. Chem. 43 (1965) 2530. m C. B. Monk, / . Amer. Chem. Soc. 70 (1948) 3281. n M. Spiro, / . Phys. Chem. 60 (1956) 976.

1 y 97-8 ±1-5

1-78 ± 0 0 3 1-85 ± 0 0 3

1-83 ± 0 0 2

I-O, 1-786 — 1-827 I-O(H), 1-898 — 1-939

6

7

8

6 7

X-ray

X-ray

X-ray

X-ray

X-ray

X-ray Neutron

NH4IO3

Ced0 3 ) 4 ,H20·

Ca(I0 3 ) 2 ,6H 2 0

Sra0 3 ) 2 ,H 2 0

Zr(I0 3 ) 4

KIO3, HIO3

1-79 ± 0 0 1

1-90 ±003

1-80 J

J

Ί ί 97-1+0-3 f

1-765 ± 0 0 0 8 1-806 + 0-008 1-836 + 0 0 1 2 1-83 1-82 1-84

6

X-ray

L1IO3

6 6 6 6 6

1-67 ±002(twice) 105-7±0-7 1-54 ± 0 0 2 | 107-3 ±1-3 (twice) 1·817±0·017 98-65±0-6

3

Neutron

93-2 - 102-8

99±1 93 ± 1 (twice)

99-2±l

102-3-105-6

106-3 ±0-2

Sm(Br0 3 )3,9H 2 0

1-485 ± 0 0 0 4

3

(°)

zo-x-o

Neutron

106

1

Ba(C10 3 ) 2 ,H 2 0

1-45 ± 0 0 3

KX-O) (A)

Intraionic dimensions

X-ray

Coordination number of halogen

NH4CIO3

X-ray or neutron diffraction

3

Compound

317±001 2-85 ± 0 0 1 (twice) 3-22 ± 0 0 1 2-55 ± 0 0 2 2-83 ± 0 0 2 2 - 9 4 ± 0 0 2 (twice) 311 ± 0 0 2 | 2-468 - 3-423

j

2-93, 2-99, 3 0 0 2-56, 2-78, 2-99 2-51, 2-73 2-55, 2-66 2-85±0-02

L

2-778, 2-819, 2-830

2-873±0-017

Intermolecular r(I · · · OI0 2 ) (Ä) R. B. Gillespie, P. K. Gantzel and K. N . Trueblood, Acta Cryst. 15 (1962) 1271. S. K. Sikka, S. N . Momin, H. Rajagopal and R. Chidambaram, / . Chem. Phys. 48 (1968) 1883. S. K. Sikka, Acta Cryst. Α25 (1969)621. A. Rosenzweig and B. Morosin, Acta Cryst. 20 (1966) 758. E. T. Keve, S. C. Abrahams and J. L. Bernstein, / . CÄem.PA^.54(1971)2556 J. A. Ibers, Acta Cryst. 9 (1956) 225.

Reference

L. Y. Y. Chan and F. W. B. Einstein, Canad. J. Chem. 49(1971)468;G.Kemper, A. Vos and H. M. Rietveld, ibid. 50 (1972) 1134.

310 j 3 00 2-86±0-07 A. Braibanti, A. M. M. Lan2-89±0-03 fredi, M. A. Pellinghelli and A. Tiripicchio, Inorg. Chim. Acta, 5 (1971) 590. A. M. M. Lanfredi, M. A. Pellinghelli, A. Tiripicchio and M. T. Camellini, Acta Cryst. B28 (1972) 679. A. C. Larson and D . T. Cromer, Acta Cryst. 14 (1961) 128.

KI · · OH2) (A)

I - O distances

TABLE 70. AVERAGE INTERATOMIC DISTANCES AND INTERBOND ANGLES IN CRYSTALLINE CHLORATES, BROMATES AND IODATES

Other studies: NaC10 3 bc ; KCKV; NaBr0 3 d ; KBr0 3 e ; NH 4 Br0 3 e ; Zn(Br0 3 ) 2 6H 2 O f ; Hg(OH)Br0 3 e ; Nd(Br0 3 ) 3 9H 2 O h ; NaI0 3 j ; Cu(0H)I0 3 j ; Ce(I0 3 ) 4 k ; K 2 H(I0 3 ) 2 CP; NH 4 CrI0 6 m . a Contains four unique iodine atoms. b W. H. Zachariasen, Z. Krist. 71 (1929) 517. c J. G. Bower, R. A. Sparks and K. N. Trueblood, U.S. Govt. Research Rept. 32 (1959) 119. d W. H. Zachariasen, Norske Videnskapsselsk Skr. 4 (1928) 90. e J . E. Hamilton, Z. Krist. 100 (1938) 104; A. Santoro and S. Siegel, Acta, Cryst. 13 (1960) 1017. ' S . H. Yü and C. A. Beevers, Z. Krist. 95 (1936) 426. * G. Bjornlund, Acta Chem. Scand. 25 (1971) 1645. h L. Helmholz,/. Amer. Chem. Soc. 61 (1939) 1544. 1 C. H. MacGillavry and C. L. P. van Eck, Rec. Trav. Chim. 62 (1943) 729. i S. Ghose, Acta Cryst. 15 (1962) 1105. k D. T. Cromer and A. C. Larson, Acta Cryst. 9 (1956) 1015. 1 A. Braibanti, A. Tiripicchio and A. M. M. Lanfredi, Chem. Comm. (1967) 1128. m K.-A. Wilhelmi and P. Löfgren, Acta Chem. Scand. 15 (1961) 1413.

TABLE 70 (cont.)

OXYACIDS AND OXYSALTS OF THE HALOGENS

1427

TABLE 71. SPECTROSCOPIC AND OTHER INVESTIGATIONS OF CRYSTALLINE CHLORATES, BROMATES AND IODATES

Measurements Spectroscopic measurements Infrared spectra Raman scattering Nuclear quadrupole resonance Mössbauer spectrum X-ray spectra emission E.S.C.A. fluorescence Thermal decomposition Calorimetry Thermal gravimetric analysis; differential thermal analysis Other thermochemical studies

Chlorates

Bromates

Iodates

1-6 1,4,16 18-21 (35C1)

2, 5, 7-10 16 22-24

2, 11-15 11,17 25

26

27,28

(79ßr, 8ißr)

29 30 32

(1271)

31

33-35

35

36-39 43

39-42 43

39 43-45

1 Infrared (absorption and reflection) and Raman study of NaC103 single crystal; J. L. Hollenberg and D. A. Dows, Spectrochim. Acta, 16 (1960) 1155. 2 Infrared spectra of NaC103, KCIO3, Ba(C103)2,H20, NaBr0 3 , KBr0 3 , AgBr03, NaI0 3 , KI0 3 , Ca(I0 3 ) 2 ,6H 2 0; F. A. Miller and C. H. Wilkins, Anal. Chem. 24 (1952) 1253; F. A. Miller, G. L. Carlson, F. F. Bentley and W. H. Jones, Spectrochim. Acta, 16 (1960) 135. 3 Infrared spectra of alkali halides doped with C103""; W. E. Klee,Z. anorg. Chem. 370 (1969) 1; G. N. Krynauw and C. J. H. Schutte, Z.phys. Chem., N.F. 55 (1967) 8, 121. 4 Infrared and Raman spectra of crystalline NaC103; Raman spectra of molten LiC103, NaC103, KC103 and AgC103; D. W. James and W. H. Leong, Austral. J. Chem. 23 (1970) 1087. 5 Infrared reflection spectra of NaC103 and NaBr0 3 single crystals; R. Duverney, Compt. rend. 254 (1962) 1954; M. Debeau, ibid. 256 (1963) 109. 6 Infrared spectra of molten alkali-metal chlorates; J. K. Wilmshurst, / . Chem. Phys. 36 (1962) 2415. 7 Infrared spectra of bromates of Na, K, Ag, Ba, Sr, Pb, Co, Ni, Mg, Zn and Cu; C. Rocchiccioli, Compt. rend. 249 (1959) 236. 8 Infrared spectra of isotopically labelled 0<*O, "O) CsBr0 3 ; C. Campbell and J. J. Turner, / . Chem. Soc. (A) (1967) 1241. 9 Infrared spectra of rare-earth bromates; G. M. Yakunina, S. E. Kharzeeva and V. V. Serebrennikov, Russ. J. Inorg. Chem. 14 (1969) 1541. !0 Infrared reflection spectrum of NaBr0 3 ; R. Duverney, C. Deloupy and R. Lalauze, Compt. rend. 260 (1965) 5749; M. Galtier, J. Barcelo and C. Deloupy, ibid. 265B (1967) 1322. 11 Infrared and Raman spectra of NaI0 3 ; P. M. A. Sherwood and J. J. Turner, Spectrochim. Acta, 26A (1970) 1975. 12 Infrared spectrum of NH 4 I0 3 ; P. W. Schenk and D. Gerlatzek, Z. Chem. 10 (1970) 153. 13 Infrared spectra of 20 iodates and iodato-complexes, including KH(I0 3 ) 2 ; W. E. Dasent and T. C. Waddington, / . Chem. Soc. (1960) 2429. 14 Infrared spectra of cobalt iodates and iodato-complexes; H. E. LeMay, jun., and J. C. Bailar, jun., / . Amer. Chem. Soc. 89 (1967) 5577. !5 Infrared spectra of polyiodates; T. Dupuis, Mikrochim. Acta (1962) 289. ™ Raman spectra of NaC103 and NaBr0 3 ; J. P. Mathieu, / . Chim. Phys. 46 (1949) 58. M Raman spectrum of NaI0 3 ; J. R. Durig, O. D. Bonner and W. H. Breazeale, / . Phys. Chem. 69 (1965) 3886. !8 35C1 nqr spectra of 16 chlorates; R. W. Moshier, Inorg. Chem. 3 (1964) 199.

1428

CHLORINE, BROMINE, IODINE AND ASTATINE: A. J. DOWNS AND C. J. ADAMS

Table 71 (cont.) 19 35C1 nqr spectrum of N a C 1 0 3 ; T. Wang, Phys. Rev. 9 9 (1955) 566. 20 35C1 nqr spectrum of KCIO3; T. Kushida, G. B . Benedek and N . Bloembergen, Phys. Rev. 104 (1956) 1364; D . B. U t t o n , / . Chem. Phys. 47 (1967) 371. 2135Q nqr spectrum of A g C 1 0 3 ; L. Anderson, Bull. Amer. Phys. Soc. 5 (1960) 412. 22 79ßr nqr spectra of N a B r 0 3 and K B r 0 3 ; R. F . Tipsword, J. T. Allender, E. A . Stahl, jun. and C. D . Williams, / . Chem. Phys. 4 9 (1968) 2464. 23 79ßr nqr spectra of bromates of Li, N a , K, A g , Ca, Sr and B a ; K. Shimomura, T. Kushida and N . Inoue, / . Chem. Phys. 2 2 (1954) 350. 24 8ißr nqr spectra of N a B r 0 3 single crystal; R. Fusaro and J. W . D o a n e , / . Chem. Phys. 4Π (1967) 5446. 251271 nqr spectra of K I 0 3 and C a ( I 0 3 ) 2 , 6 H 2 0 ; G. W. Ludwig, / . Chem. Phys. 25 (1956) 159. 26 83ßr Mössbauer spectrum of K B r 0 3 ; M . Pasternak and T- Sonnino, PÄy,s. Rev. 164 (1967) 384. 27 1271 Mössbauer spectra (ZnTe source) of N a I 0 3 and N a 3 M n ( I 0 3 ) 6 ; P. Jung and W. Triftshäuser, Phys. Rev. 175 (1968) 512. 281291 Mössbauer spectra of K I 0 3 , N H 4 I 0 3 and B a ( 1 0 3 ) 2 ; D . W. Hafemeister, G. de Pasquali and H . de Waard, Phys. Rev. 135B (1964) 1089. 29 χ - r a y emission spectrum of N a C 1 0 3 ; V. I. Nefedov, / . Struct. Chem. 8 (1967) 919; W. Nefedow, Phys. Stat. Solidi, 2 (1962) 9 0 4 . 30 N a C 1 0 3 ; A . Fahlman, R. Carlsson and K. Siegbahn, Arkiv. Kemi, 2 5 (1966) 301. 3i K I 0 3 ; C. S. Fadley, S. B. M. Hagstrom, M . P. Klein and D . A . Shirley, / . Chem. Phys. 48 (1968) 3779. 32 Spectrum of N a C 1 0 3 ; J. A . Bearden, Rev. Mod. Phys. 39 (1967) 86; D . S. Urch, / . Chem. Soc. (A) (1969) 3026. 33 K C 1 0 3 decomposition; A . A . Gilliland and D . D . Wagman, / . Res. Nat. Bur.

Stand. 69\ (1965)1.

3 4 N a C 1 0 3 and K C 1 0 3 decomposition; A . F . Vorob'ev, N . M . Privalova, A . S. Monaenkova and S. M . Skuratov, Doklad. Akad. Nauk S.S.S.R. 135 (1960) 1388; A . F . Vorob'ev, N . M. Privalova and L. T. Huang, Vestn. Mosk. Univ. Ser. II, Khim. 18 (1963) 27. 35 K C 1 0 3 and K B r 0 3 decomposition; G. K. Johnson, P. N . Smith, E . H . Appelman and W. N . Hubbard, Inorg. Chem. 9 (1970) 119. 36 Decomposition of alkali-metal chlorates; M . M. Markowitz, D . A . Boryta and H . Stewart, jun., / . Phys. Chem. 6 8 (.1964) 2282. 37 A g C 1 0 3 decomposition; F . Solymosi, Z. phys. Chem., N.F. 57 (1968) 1. 3 8 N H 4 C 1 0 3 decomposition; F . Solymosi and T. Bansagi, Combust. Flame, 13 (1969)262. 39 Decomposition of chlorates, bromates and iodates of K, R b and C s ; O. N . Breusov, N . I. Kashina and T. V. Revzina, Russ. J. Inorg. Chem. 15 (1970) 316. 40 Decomposition of 17 alkali-metal, alkaline earth, transition metal and rare earth bromates; G. M . Bancroft and H . D . G e s s e r , / . Inorg. Nuclear Chem. 2 7 (1965) 1545. 41 Decomposition of hydrated bromates of divalent metals; B. Du§ek and F . Petra, Russ. J. Inorg. Chem. 13 (1968) 307. 42 Decomposition of hydrated bromates of rare-earth elements; I. Mayer and Y . Glasner, / . Inorg. Nuclear Chem. 2 9 (1967) 1605. 43 Calculation of lattice energies; A . Finch and P. J. Gardner, / . Phys. Chem. 69 (1965) 384. 44 Reaction of K I 0 3 ( c ) with aqueous hydrazine; P. B. Howard and H . A . Skinner, / . Chem. Soc. (A) (1967) 269. 4 * Reaction of K I 0 3 ( c ) with aqueous H I ; C . W u , M . M . Birky and L. G. Hepler, / . Phys. Chem. 67 (1963) 1202.

OXYACIDS AND OXYSALTS OF THE HALOGENS

1429

TABLE 72. HALOGEN ATOM CHARGES IN CHLORATE, BROMATE AND IODATE IONS

Charge on halogen atom in Technique Infrared intensity measurements Raman intensity measurements Mössbauer spectroscopy X-ray photoelectron spectroscopy Nuclear quadrupole resonance spectroscopy6 X-ray emission spectroscopy

cio 3 -

Br0 3 -

a ca. + 1*

+ 1-72* ca. + 11

+ 1·02* + 1-21« + 1-49* + 1-721

+0-87« +0-72«

IO3ca. + 1J +0-83° + 1-10* +0-69* +2-73«

a

Charge on oxygen estimated to be — 0·25; chlorine charge not estimated; G. N. Krynauw and C. J. H. Schutte, Z. phys. Chem., N.F. 55 (1967) 121. to R. Duverney, C. Deloupy and R. Lalauze, Compt. rend, 260 (1965) 5749. c D. W. Hafemeister, G. de Pasquali and H. de Waard, Phys, Rev. 135B (1964) 1089. d C. S. Fadley, S. B. M. Hagstrom, M. P. Klein and D. A. Shirley, / . Chem. Phys. 48 (1968) 3779. e E. A. C. Lücken, Nuclear Quadrupole Coupling Constants, p. 285, Academic Press (1969). f Calculated assuming tetrahedral bond angles. * Calculated using experimental bond angles. * R. Manne, / . Chem. Phys. 46 (1967) 4645; A. Fahlman, R. Carlsson and K. Siegbahn, Arkiv. Kemi, 25 (1966) 301. 1 W. Nefedow, Phys. Stat. Solidi, 2 (1962) 904. *G. W. Chantry and R. A. Plane, / . Chem. Phys. 32 (1960) 319.

The results of various spectroscopic investigations have been used to calculate atomic charges for the halate ions; some of these values are reproduced in Table 72. In view of the assumptions necessary in analysing the data, the discrepancies between the different values are hardly surprising; although all the calculations imply considerable polarity in the X-O bond, there is no hint of any trend in the derived charges. Reactions Decomposition Reactions Dilute solutions of HCIO3 and HBr0 3 decompose if heated, but the cold solutions may be concentrated in vacuo. While the decomposition of bromic acid is described by the equation 4HBr0 3 -► 2Br 2 +50 2 +2H 2 0

chloric acid disproportionates, possibly with explosive violence. Dilute solutions evolve chlorine and oxygen in a reaction approximating to 8HCIO3 -> 4HC10 4 +2H 2 0+302+2Cl 2

but in more concentrated solution chlorine dioxide is one of the products. Large amounts of the dioxide result from the action of concentrated sulphuric acid on solid chlorates (p. 1367). 3HC103 -> HCIO4+H2O+2CIO2

1430

CHLORINE, BROMINE, IODINE AND ASTATINE: A. J. DOWNS AND C. J. ADAMS

Three decomposition processes are possible for the halate ions: (1) (2) (3)

4X03-->X-+3X042X03-->2X-+302 4X
At temperatures near to 25°C reaction (1) is thermodynamically feasible only for chlorates, as witnessed by the thermodynamic data for the potassium salts (Table 73) and by the equilibrium constants in aqueous solution (CIO3 ~, 1022; Br0 3 ~, 10 - 3 0 ; IO3-, 10 ~ 53 ). Disproportionation proceeds very slowly even in boiling chlorate solutions; it is observed, however, as an exothermic process in the pyrolysis of metal chlorates. For the molten alkali-metal salts reaction (1) comprises the major decomposition route at temperatures below 500°C, the melt becoming viscid as perchlorate is formed; at higher temperatures the perchlorates are themselves unstable. For other chlorates, though observed initially, dis­ proportionate is superseded by other decomposition processes. While there is no evidence for formation of BrC>4 ~ or IO4 ~ during the pyrolysis of bromates or iodates, the anhydrous alkaline earth iodates decompose at high temperatures according to the reaction 5M(I0 3 ) 2 -> M 5 ( I 0 6 ) 2 + 4 I 2 + 9 0 2

Perhalate formation is supposedly an internal oxygenation process. Disproportionation apart, whether decomposition of metal halates occurs by reaction (2) (with formation of the metal halide) or by reaction (3) (with oxide-formation) is deter­ mined mainly on thermodynamic grounds, by the relative stabilities of the metal halide and oxide, but also kinetically, by the relative activation energies of the two processes723; on both counts oxide-formation is favoured by the presence of a strongly polarizing cation. The effects of particle size and rate of heating are also important kinetically. Halide-formation is observed for alkali-metal, alkaline earth and silver salts, while magnesium, transition metal and rare earth salts favour oxide-formation, occasionally with explosive evolution of chlorine dioxide from chlorates, e.g. from Cu(C103)2. Barium bromite may be prepared by the controlled pyrolysis of barium bromate699. The thermal stability of the salts falls in the sequence iodate > chlorate > bromate, and also decreases as the polarizing power of the cation rises. Oxygen-evolution from alkali chlorates is catalysed by transition metal salts, which depress the temperature of decompo­ sition and preclude disproportionation. Heating mixtures of potassium chlorate and manganese dioxide has long been exploited, particularly for didactic purposes, as a method of producing oxygen on the small scale. While the initial step in the thermal decomposition of alkali-metal halates is rupture of the X-O bond, which occurs only at high temperature (> 300°C), decomposition of the ammonium halates begins at much lower temperatures (NH4CIO3, 50°C; NtLtBrC^, — 5°C; NH4IO3, ca. 100°C), and the compounds explode on further heating. It is believed that, as in the case of NH4CIO4, the reaction commences with proton transfer, NH>:(s)

► NH(s) + HXO (s)

I!

NH3(g)· 7

It

HX03(g)

23 G. M. Bancroft and H. D. Gesser, / . Inorg. Nuclear Chem. 27 (1965) 1545.

+ 17-3 + 17-7

+ 100 + 11-5

-8-4 -7-7

+ 83-8 + 69-2 d 525 (-^KI+3/2 0 2 ) 303 (->KI0 3 +l/2 0 2 )

+ 31-8 + 6-5

-25-1 -49-9

-10 -25-4

+46-5 +25-6 a 390 (->KBr+3/2 0 2 ) 275 (->KBr0 3 +l/2 0 2 )

+41-2 +22-6

-7-8 -25-9

-91 -26-8

+ 51-9 + 600 c 472 (-*3/4KC10 4 +l/4KCl) 534 (->KCl+20 2 )

-119-5 -99-6 -1101 -83-5

Iodine

-860 -64-7 -68-7 -40-7

Bromine

-951 -70-8 -103-2 -72-2

Chlorine

b c

» G. K. Johnson, P. N. Smith, E. H. Appelman and W. N. Hubbard, Inorg. Chem. 9 (1970) 119. Selected Values of Chemical Thermodynamic Properties, N.B.S. Technical Note 500, Washington (1952). M. M. Markowitz, D. A. Boryta and H. Stewart, jun.,/. Phys. Chem. 68 (1964) 2282; M. M. Markowitz, D. A. Boryta and R. F. Harris, /. Phys. Chem. 65(1961)261. d O. N. Breusov, N. I. Kashina and T. V. Revzina, Russ. J. Inorg. Chem. 15 (1970) 316.

κχο3 κχο4

Ajyr0(c), KX03(kcal mol-*)* AG,°(c), KX03(kcal mol"*)· MIf°(c), KX04(kcal mol"*)» AG,°(c), KX04(kcal mol-i)* KX03(c) -* KX(c)+3/2 02(g)»'b Atf°(kcalmol-i) AG°(kcalmori) KX04(c) -> KX(c)+2 02(g)·1» Ajy°(kcalmol-i) AG°(kcalmol-i) KX03(c) -► 3/4 KX04(c)+1/4 KXfc)·.* Aif0(kcalmol-i) AG°(kcalmol-i) KXOÄ(s) -> 1/2 K20(s)+1/2 X 2 (g)+(2/t- 1)/4 02(g)».* /i = 3;Ajr(kcalmol-i) n = 4;A^°(kcalmol-i) Decomposition temperature (°C)

x=

TABLE 73. THERMODYNAMIC PROPERTIES OF POTASSIUM HALATES AND PERHALATES AT 25°C

1432

CHLORINE, BROMINE, IODINE AND ASTATINE: A. J. DOWNS AND C. J. ADAMS

followed by rapid decomposition of the acid and oxidation of ammonia to give the products, which include a residue of NH4NO3. Flash photolysis of the XO3 - anions632 in aqueous solution (Table 54) produces X 0 2 radicals by the reaction XO3 " ,H 2 0 - * (XO3 - ,H 2 0)* -► X 0 2 + O H 4- OH ~

Depending on the temperature of irradiation and subsequent annealing processes, the action on crystalline chlorates610*724 and bromates603'725 of X-rays or y-rays can produce many species, including perhalate ions, oxyhalogen radicals (Table 54) and radical anions (p. 1460). XO3-

X or y

► XO4-, X 0 2 - , X O - , X - , XO3, X 0 2 , X-XO3-, XO3 2 -, 0 2 and O3 (X = Cl or Br)

Similar decay follows the capture of thermal neutrons by the halates726. Oxidation by Chlorate, Bromate and Iodate Both thermodynamically and kinetically the oxidizing powers of the halates in aqueous solution are marked functions of the hydrogen ion concentration. The potentials for the conversion of halate to halide (the usual reduction process) are more favourable in acid solution (1-1-1-5 volts as against 0-3-0-7 volt in alkaline solution), and the reactions are more rapid. In alkaline media chlorate exhibits no oxidizing capabilities, and its reduction in weakly acidic solution may be very slow, requiring catalysis by transition metal deriva­ tives, e.g. of Os(VIII), V(V) or Mn(VII). In concentrated solution chlorates are often reduced to chlorine dioxide rather than to chloride. Kinetic differences apart, the three halate anions are quite similar in their oxidizing properties; some reactions of bromates are summarized in Scheme 11. Bromates and iodates undergo some redox reactions in alkaline solution; thus, reaction with hydroxylamine or hydrazine (as their sulphates) produces nitrogen (alkaline chlorates not reacting), and iodate oxidizes vanadyl salts to the pentavalent state: 6V02 + +I0 3 - + 180H- - > 6 V 0 3 - + I - + 9 H 2 0

However, many transformations which are rapidly wrought by the alkaline halites or hypohalites at room temperature, e.g. the oxidation of As(III) to As(V), take place only slowly with the halates, even in acid solution. Halates and halides react in acid solution in all nine possible combinations: iodides are oxidized quantitatively to iodine, and bromides to bromine (setting aside subsequent polyhalide-formation): the liberation of iodine is a common procedure in analysing halates. XO3- + 5X- +6H + X0 3 - + 6Y-+6H +

-> 3X 2 +3H 2 0 (X = Cl, Br or I) -^X" + 3Y 2 +3H 2 0 (X = C1; Y = IorBr; X = Br; Y = I) XO3- + 5Y- +6H + -> 2Y 2 +XY+3H 2 0 (X = I; Y = Br) 2C10 3 -+2C1-+4H + ->C1 2 +2C10 2 +2H 2 0

Chlorides produce both Cl2 and CIO2 with chlorate, bromine and chlorine with bromate, and chlorine and iodine chlorides with iodate. The reactions share common kinetic features with the (halide-catalysed) exchange of oxygen between the halate ions and water, which is 724 H . G. Heal, Canad. J. Chem. 37 (1959) 979. 725 L . C. Brown, G. M. Begun and G. E. Boyd, / . Amer. Chem. Soc. 91 (1969) 2250. 726 T . Andersen, H. E. L. Madsen and K. Olesen, Trans. Faraday Soc. 62 (1966) 2409.

1433

OXYACIDS AND OXYSALTS OF THE HALOGENS

I2,IO~ (excess Br0 3 )

CrO"

SCHEME 11. Reactions of aqueous bromates.

also acid-catalysed. The general rate equations have theform™,727,728 rate =

Ar[X0 3 -][H + P[Y-]

and the proposed mechanism involves an intermediate of the type YXO2: XO3- + 2 H + ^ M ± r [H2X03+] + Y-

^H2ö

[H2XOj]

+ [YXO^-JCY,

Η,Ο

+

XO-

γ χ ο 2 (X,Y =*= ci)

r

[H2xo3 ] + Y~ Y2 +

2XO,

Subsequent reaction of XO2 ~ with Y ~ is assumed to be rapid. It is uncertain whether the YX0 2 intermediate has the structure Y - X \

or O-X-O-Y. The former is favoured by

comparison with known compounds (the halogenylfluorides),but the latter structure could also account for the kinetically related oxidation of NO2 ~ and SO32 ~, which involves quanti­ tative transfer of oxygen from the halate ion to the substrate. Halide ion catalysis has been found for other oxidation reactions of the halate ions. Iodine replaces chlorine and bromine in CIO3"" and BrOß": l2+2X0 3 -->X2+2I0 3 727 A . F . M. Barton and G. A. Wright, J. Chem. Soc. (A) (1968) 1747. 728 M. Anbar and S. Guttmann, / . Amer. Chem. Soc. 83 (1961) 4741.

1434

CHLORINE, BROMINE, IODINE AND ASTATINE: A. J. DOWNS AND C. J. ADAMS

The bromate-iodine reaction has the following stages at pH 1-5-2-5 729 : (1) (2) (3) (4)

A n induction period in which a catalyst is produced (probably HOBr) I2 + B r 0 3 - - + I B r + I 0 3 3IBr+2Br0 3 " +3H 2 0->5Br" +3I0 3 " +6H + 5Br " + Br0 3 " + 6H + -> 3 B r 2 + 3 H 2 0

The reaction between iodate and sulphite in acid solution (Landolt's reaction) is interesting: three different steps are involved which give rise to a periodic appearance and disappearance of iodine in solution. (1) I0 3 -+3S0 3 2- ->I +3S04 2 (2) 5 I - + I 0 3 - + 6 H + - + 3 I 2 + 3 H 2 0 (3) 3I 2 +3S0 3 2" +3H 2 0 -> 61" +6H + + 3S0 4 2 "

The oxidation of HN 3 by Br0 3 ~ in perchloric acid solution affords N 2 , N 2 0, OBr ~ and Br2; the oxygen in N 2 0 is derived mostly from the solvent730. Solid chlorates readily oxidize organic material and non-metals such as sulphur, selenium, tellurium, phosphorus and arsenic. Dry mixtures are generally stable, but traces of moisture may induce ignition or explosive decomposition (via the formation of HCIO3). Bromates and iodates also form unstable mixtures with combustible or oxidizable sub­ stances. Molten chlorates have been used to oxidize iodide to iodate. Oxidation of Halates Oxidation of halogen(V) to halogen(VII) is thermodynamically more facile in alkaline solution than in acid solution. Perchlorates arise in the disproportionation of chlorates, from which they may also be obtained by electrolytic or chemical oxidation, e.g. with persulphate (E° = +2-01 volts). Iodates are more readily oxidized, since chlorine (E° = +1-36 volts), bromine (+1-07 volts), hypochlorite (+0-89 volt), permanganate (+1-23 volts) or per­ sulphate effect the change; anodic oxidation is hindered by reduction of iodate at the cathode. The problem of oxidizing bromates has perplexed chemists for many years. Though the synthesis of perbromates has now been accomplished663»731 using F2, XeF2 or electrical means to oxidize aqueous Br0 3 -, the reduction potential of the Br0 4 ~/Br03 _ couple (+1 -76 volts in acid solution), though greater than those for the other halogen(VII)/halogen(V) couples, is not sufficiently great to explain why ozone or persulphate does not effect the oxidation. A detailed kinetic study has been made of the oxidation of halates by XeF2 in aqueous solution732; the optimum yields are CIO4 -, 93%; Br0 4 -, 12%; I 0 4 -, 93%. The oxidation of C103 ~ and Br0 3 - is accomplished by an intermediate in the reaction of XeF2 with water, producing a second intermediate which either goes on to give perhalate or goes back to halate; the latter reaction predominates for bromate. In dilute solution iodate-oxidation entails a similar mechanism, but in concentrated iodate solutions there is a direct reaction via a xenon iodate: XeF 2 +I0 3 "

> [FXeOI02] —%—+ Xe+HF+H + +IO4-

Fluorination of solid potassium chlorate affords C103F (p. 1391). 729 D . E. C. King and M . W. Lister, Canad. J. Chem. 46 (1968) 279. 730 R . c . Thompson, Inorg. Chem. 8 (1969) 1891. 731 E. H . Appelman, Inorg, Chem. 8 (1969) 223. 732 E . H . Appelman, Inorg. Chem. 10 (1971) 1881.

OXYACIDS AND OXYSALTS OF THE HALOGENS

1435

Uses Sodium chlorate is manufactured in very large amounts for conversion into (i) chlorine dioxide (by reduction with chloride or sulphur dioxide), and (ii) perchloric acid and perchlorates (by electrolysis in acid solution). Quantities are also used as oxidizers in the making of matches and in fireworks, and form a common ingredient of weed-killing preparations. The commercial applications of bromates and iodates depend on their ability to oxidize thiolic residues in proteins to S-S linkages, a characteristic which is utilized in improving the baking properties of flours, and in preparations for waving hair. Potassium bromate and iodate are important primary analytical standards733 ; bromates are used especially in the determination of trivalent arsenic and antimony, by oxidation from the trivalent to the pentavalent state, while iodate is used in numerous procedures for the estimation of a variety of metals. (E) PERHALIC ACIDS AND PERHALATES

Introduction Discovery Potassium perchlorate was discovered in 1816 by von Stadion, who treated fused potas­ sium chlorate with sulphuric acid; chlorine dioxide was evolved and potassium perchlorate was crystallized from the mixture. Subsequent investigations by the same worker led to the isolation of perchloric acid as a result of distilling KCIO3-H2SO4 mixtures, and to the electrolytic oxidation of KCIO3 in aqueous solution. The nature of the disproportionation of chlorates was first truly recognized by Serullas (in 1830) when he isolated KCIO4 from pyrolysed KCIO3 without recourse to treatment with sulphuric acid; he also prepared perchloric acid solutions by the reaction 2KC104+H2SiF6 -> K2SiF6 ! +2HC104 Perchlorates have since been found in natural nitrate and saline deposits, and in sea water. Interest in perchlorates waned until about 1894, when manufacture by electrolysis of chlorate solutions was undertaken in Sweden for use in explosives. Since that date, the level of commercial production of perchlorate has fluctuated in inverse proportion to the extent of world peace; current world production is in excess of 50,000 tons per annum, much of which is put to use as oxidizers in rocket fuels. The natural interest of national governments in these applications has meant that the physicochemical properties of perchloric acid and its derivatives have been investigated in great detail. There have been several comprehensive reviews of the chemistry of perchloric acid and the perchlorates572»575b,576,734-738; two of these articles give special emphasis to the physical chemistry of perchloric acid736»73?. 733 j . H . Thompson, Comprehensive Analytical Chemistry (ed. C. L. Wilson and D . W. Wilson), Vol. IB, p. 259, Elsevier (1960). 734 J. C. Schumacher (ed.), Perchlorates, A.C.S. Monograph No. 146, Reinhold, New York (1960). 735 j . c . Schumacher and R. D. Stewart, Kirk-Othmer Encyclopedia of Chemical Technology, 2nd edn., Vol. 5, pp. 61-84, Interscience, New York (1964). 736 G. S. Pearson, Adv. Inorg. Chem. Radiochem. 8 (1966) 177. 737 A. A. Zinov'ev, Russ. Chem. Rev. 32 (1963) 268. 738 c . C. Addison, Mellor's Comprehensive Treatise on Inorganic and Theoretical Chemistry, Supplement II, Part I, pp. 596-620, Longmans, London (1956).

1436

CHLORINE, BROMINE, IODINE AND ASTATINE: A. J. DOWNS AND C. J. ADAMS

Periodic acid and its salts575**'576'624»739»740 have been known for many years; thus, the oxidation of iodates by chlorine was reported as early as 1833. By contrast, perbromates were not discovered until 1968 during studies of the /?-decay of 82SeC>42~, an event rapidly followed by the disclosure of chemical syntheses of perbromates. The difficulties encountered in oxidizing bromates were described in the preceding section, although the inaccessibility of perbromates is due not so much to their intrinsic properties as to the high-energy kinetic barrier which opposes the conversion of Br(V) to Br(VII). A Note of Caution It must be emphasized that, despite their kinetic inertness at ambient temperatures, perchloric acid and the perchlorates are potentially very strong oxidizing agents. Anhydrous perchloric acid is subject to spontaneous explosive decomposition on storage or in the presence of contaminants, and many of the perchlorates have highly exothermic decompo­ sition processes open to them, e.g. AgC10 4 -> AgCl+ 2 0 2

ΔΗ

25 kcal mol" i

which bring considerable risks of explosion to operations such as drying and grinding, especially if the cation is readily oxidizable. Though 72% perchloric acid is an extremely useful and safe analytical reagent, the dangers involved in mixing 100% perchloric acid with organic material would be hard to exaggerate: in 1947 an explosion caused by the introduction of some plastic into a tank containing 220 U.S. gallons of a mixture of per­ chloric acid and acetic acid resulted in 17 deaths and the destruction of 116 buildings. Organic salts of perchloric acid have a bad record of laboratory explosions; arene diazonium perchlorates were once thought to be the most explosive compounds known. The prepara­ tion of anhydrous perchloric acid may be prohibited by local regulations, and should in any case be undertaken only in small amounts; the literature contains detailed accounts of appropriate safety precautions571»734. Comparison of Perchlorates, Perbromates and Periodates Some aspects of the chemistry of the perhalic acids and their derivatives are summarized in Table 74. Though data for the perbromates are still sparse, recent results show that they have much in common with the perchlorates, whereas the chemistry of the periodates is rather different and resembles that of the tellurates(VI). Many of the differences between periodates and the corresponding chlorine and bromine derivatives are attributable, at least in part, to the greater size of the iodine atom. While the perchlorates and perbromates are found only with tetra-coordinated halogen atoms, hexa-coordination about iodine is the rule rather than the exception in periodates. A much-quoted example concerns the phases having the composition H 5 X0 6 , where X = Cl or I; these are crystalline at room temperature with the structure [H5O2] +C1C>4~ in the one case or (HO)5IO in the other, whereas anhydrous HCIO4 is liquid under similar conditions. HIO4 may be made by controlled dehydration of (HO)5IO, and IO4- is the principal con­ stituent of alkaline periodate solutions; the anion enters into a rapid hydration-dehydration equilibrium in solution: IO4- + 2 H 2 0 <± (HO) 4 I0 2 739 G. J. Hills, Mellor's Comprehensive Treatise on Inorganic and Theoretical Chemistry, Supplement II, Part I, pp. 896-907, Longmans, London (1956).

740 H . Siebert, Fortschritte

Chem. Forsch, 8 (1967) 470.

OXYACIDS AND OXYSALTS OF THE HALOGENS

1437

TABLE 74. GENERAL COMPARISON OF THE HALOGEN(VII) OXYACIDS

Halogen Property

Chlorine

Structure of the acid

HOCIO3

Physical state at 20°C Acid character

liquid very strong, monobasic

Anions found in solution

Oxidizing character of aqueous solution

Oxygen-exchange between H2O and XO4-

Bromine HOBr0 3 probably

—

CKV

very strong, monobasic Br0 4 ~

absent at 20°C

sluggish at 20°C

strong at 100°C

strong at 100°C

not observed

not observed

Iodine (HO)5IO solid weak, polybasic H4I06H 3 I0 6 2H 2 I0 6 3H2I2O104IO4potent and rapid at 20°C

extremely rapid

which is not observed for the other perhalates. The hydrated periodates persist in acid solution and are isolable as salts. Periodates also form polymeric anions with I-O-I bridges, as well as very stable complexes and heteropolyanions with I-O-M bridges; no analogous polymerization of perchlorates (except to give the anhydride CI2O7) or perbromates has been reported, and perchlorato-complexes contain only weakly bound CIO4" ions. The lability of the groups attached to iodine is also responsible for the rapidity with which periodates oxidize many materials. While perchlorates and especially perbromates are also strong oxidizing agents, their reactions are hampered at room temperature by the absence of suitable intermediates of low energy; nucleophilic attack on the small halogen atom to give afive-coordinateintermediate is difficult because of the relatively congested tetrahedron of negatively charged oxygen atoms surrounding it, while the alternative path­ way—rupture of a Cl-0 or Br-O bond—is also expensive of energy. Simple perchlorates and perbromates ΜχΧθ4, like many other salts containing tetrahedral anions, crystallize at room temperature with the orthorhombic Barite (BaSC>4) structure; at higher temperatures (>200°C) perchlorates transform to a cubic lattice, a change associated with the onset of free rotation of the CIO4 - ions. Metaperiodates M x I0 4 adopt the tetragonal Scheelite structure. The anions XO4- have been characterized by spectroscopic and single-crystal diffractometric studies (Tables 75, 76, 79 and 83); in simple salts the anions are tetrahedral within the limits of accuracy of X-ray diffraction ex­ periments, although the vibrational spectra of the solids show effects, such as the appearance of a weak infrared absorption due to the formally forbidden symmetric stretching mode vx(ai), arising out of the occupation by the anion of a site of symmetry lower than Td. The bond lengths increase from ca. 1-45 A in perchlorates to 1-61 A in KB1O4 and 1-78 A in NaI04. While the valence force constants decrease from 8-2 mdyne A" 1 for CIO4-to ca. 6 0 mdyne A - 1 for Br0 4 - and I 0 4 ~, the decreases in the deformation and interaction force

1438

CHLORINE, BROMINE, IODINE AND ASTATINE: A. J. DOWNS AND C. J. ADAMS TABLE 75. PROPERTIES OF THE AQUEOUS PERHALATE IONS Χθ4~

Thermodynamic functions at 298°K Af/>°(kcalmol-i) a AG/°(kcalmol-i) a S^caldeg-imol- 1 )» Reduction potentials Acid solution £°(X0 4 -/X03-)(volts) a £ 0 (X0 4 -/£X 2 )(volts) a · 0 E 0 (X0 4 -/X-)(volts) a · 0 Alkaline solution £°(X04-/X0 3 -)(volts) a . c £°(X0 4 -/iX2)(volts) a ' c £°(X0 4 -/X-)(volts) a . c Spectroscopic properties Raman spectrum Infrared spectrum Fundamental frequencies (cm - 1 ) "i(ai)(R) v 2 (e)(R) v 3 (/ 2 )(R,i.r.) v 4 (/ 2 )(R,i.r.) Force constants (mdyne A" 1 )* fr fei* feelr2 (frt-MIr Ultraviolet-visible absorption spectrum Amax(nm) € max(M"l)

{

cio 4 -

BrCXr

io4-

-30-70 -1-84 43-6

3-19 29-18 44-7

-34-56 -1109 48-8

+ 1-23 + 1-39 + 1-39

+ 1-76 + 1-58 + 1-52

b

+0-40 +0-45 +0-56

+0-93 +0-64 +0-69

+0-81 +0-37 +0-40

d,e,f j

g,h

d,i

931 462 1107 631

801 331 878 410

791 256 853 325

8-24 0-87 -0-21 0-78 3 <185

Nmr spectrum l7 0 chemical shift (ppm with respect to H2O) - 2 8 8 ±51 35 C1 chemical shift (ppm with respect to Cl~) - 1 0 0 0 m 127 I spectrum Ionic mobility (25°C) (cm 2 o h m - 1 g-ion - 1 ) 67-4°

605 0-48 -012 0-38 h 187 3-9x103

5-90 0-30 -009 007 k 222-5 10-7x103

n 54-6p

R, Raman-active; i.r., infrared-active. a G. K. Johnson, P. N. Smith, E. H. Appelman and W. N. Hubbard, Inorg. Chem. 9(1970)119. b See Table 80 for periodate potentials in acid solution. c A. G. Sharpe, Halogen Chemistry (ed. V. Gutmann), Vol. 1, p. 13, Academic Press (1967). d Includes absolute intensity measurements; G. W. Chantry and R. A. Plane, / . Chem. Phys. 32 (1960) 319; ibid. 34 (1961) 1268. e Includes effect of metal ions on the spectrum of C10 4 ~; M. M. Jones, E. A. Jones, D. F. Harmon and R. T. Semmes,/. Amer. Chem. Soc. 83 (1961) 2038; R. E. Hester and R. A. Plane, Inorg. Chem. 3 (1964) 769. f J. W. Akitt, A. K. Covington, J. G. Freeman and T. H. Lilley, Trans. Faraday Soc. 65(1969)2701. * L. C. Brown, G. M. Begun and G. E. Boyd,/. Amer. Chem. Soc. 91 (1969) 2250. h E. H. Appelman, Inorg. Chem. 8 (1969) 223. 1 H . Siebert and G. Willfahrt, cited in H. Siebert, Fortschritte Chem. Forsch. 8 (1967) 470. J J. D. S. Goulden and D. J. Manning, Spectrochim. Actat 23A (1967) 2249. k C. E. Crouthamel, H. V. Meek, D. S. Martin and C. V. Banks, / . Amer. Chem. Soc. 71(1949)3031. 1 B. N. Figgis, R. G. Kidd and R. S. Nyholm, Proc. Roy. Soc. 269A (1962) 469.

1439

OXYACIDS AND OXYSALTS OF THE HALOGENS Table 75 (cont.) m a

Y. Saito, Canad. J. Chem. 43 (1965) 2530. R. M. Kren, H. W. Dodgen and C. J. Nyman, Inorg. Chem. 7 (1968) 446. ° J. H. Jones, / . Amer. Chem. Soc. 67 (1945) 855. p C. B. Monk, / . Amer. Chem. Soc. 70 (1948) 3281. q L. J. Heidt, G. F. Koster and A. M. Johnson, / . Amer. Chem. Soc. 80 (1959) 6471. TABLE 76. INVESTCGATIONS OF SOLID PERHALATES

Measurements X-ray crystallography Spectroscopic investigations Infrared and Raman spectra Mössbauer spectra X-ray spectra: E.S.C.A. emission fluorescence Thermal decomposition Calorimetry Thermal gravimetric analysis; differential thermal analysis

Perchlorates Perbromates

Periodates

1-7

8

9

1,10-14

15, 16

17-21 22,23

24 26 27 28 30-32

25

29 32,33

i See also Table 79. 2LiC10 4 ,r ftv (Cl-0)= 1-44 ±001 A; KCIO4, r a v (Cl-0) = l-43±002Ä; R. Prosen, Ph.D. Thesis, University of California (1955), cited in ref. 4. 3NH 4 C10 4 ,r av (Cl-0) = 1·46±0·03 A; K. Venkatesan,Proc. Indian Acad. Sei. A56 (1957) 134. 4 [N0 2 ]C10 4 ,r ÄV (Cl-0) = 1·464±0·007 A; M. R.Truter,D. W. J. Cruickshank and G. A. Jeffrey, Acta Cryst. 13 (1960) 855. 5 [N2H5]C104,*H20, r a v (Cl-0) = 1-447 ±0004 A; R. Liminga, Acta Chem. Scand. 21 (1967) 1217. 6HC104(mesitaldehyde)2, r a v (Cl-0) = 1-458±0007A; C. D. Fisher, L. H. Jensen and W. M. Schubert, / . Amer. Chem. Soc. 87 (1965) 33. 7 Cu(C6H5N3)2(C104)2, r a v ( C l - 0 ) = 1·422±0·0ΐΑ; J. J. Bonnet and Y. Jeannin, Acta Cryst. 26B (1970) 318. 8KBr0 4 ,r a v (Br~0) = 1 ·610± 0016 A; S. Siegel, B.Tani and E.H. Appelman, Inorg. Chem. 8 (1969) 1190. 9 See also Table 83. io Infrared spectra of alkali halides doped with C104~; W. E. Klee, Z. anorg. Chem. 370 (1969) 1; G. N. Krynauw and C. J. H. Schutte, Z. phys. Chem., N.F. 55 (1967) 8,121. H Infrared spectra of MC104 (M = Li, Na, K, Rb, Cs, Tl and Ag); A. Hezel and S. D. Ross, Spectrochim. Acta, 22 (1966) 1949. 12 Infrared spectra of MC104 (M = Na, K and NH4) and M(C104)2,xH20 (M = Mg, Ca, Ba, Fe, Co, Ni, Cu and Zn); S. D. Ross, Spectrochim. Acta, 18 (1962) 225. 13 Infrared spectrum of KC104 single crystal at low temperature; R. A. Schroeder, E. R. Lippincott and C. E. Weir, / . Inorg. Nuclear Chem. 28 (1966) 1397. 14 Infrared spectra of M(C104)3,«DMSO in = 7, M = Sm, Gd and Y; n = 8, M = La, Ce, Pr and Nd); Y. N. Krishnamurthy and S. Soundararajan, / . Inorg. Nuclear Chem. 29 (1967) 517; Z. anorg. Chem. 349 (1967) 220. 15 Infrared and Raman spectra of KBr0 4 ; E. H. Appelman, Inorg. Chem. 8 (1969) 223. 16 Infrared and Raman spectra of CsBr04 and Ph4AsBr04; L. C. Brown, G. M. Begun and G. E. Boyd, / . Amer. Chem. Soc. 91 (1969) 2250.

1440

CHLORINE, BROMINE, IODINE AND ASTATINE: A. J. DOWNS AND C. J. ADAMS

Table 76 (cont.) 17 See also Table 84. is Infrared and Raman spectra of M I 0 4 (M = NH 4 , Na, K, Rb, Cs and Ag); R. G. Brown, J. Denning, A. Hallett and S. D. Ross, Spectrochim. Acta9 26A (1970) 963. 19 Infrared spectra of M 5 I0 6 (M = Li and Na) and M 5 (I0 6 )2 (M = Ca, Sr and Ba); J. Hauck, Z. Naturforsch. 25b (1970) 647. 20 Infrared spectra of Na 3 H 2 I0 6 , Na 2 H 3 I0 6 , NaI0 4 ,3H 2 0, Κ4Ι2Ο9, K4l 2 09,9H 2 0, AgI0 4 , Ag 3 I0 5 , Ag 5 I0 6 , Ag 2 HI0 5 and Ag 2 H 3 I0 6 ; J. R. Kyrki, Suomen KemistilehtU B38 (1965) 192. 21 Infrared spectra of many alkali-metal periodates; H. Siebert, Z. anorg. Chem. 272 (1953) 1; ibid. 303 (1960) 162; ibid. 304(1960) 266; Fortschritte Chem. Forsch. 8 (1967) 470. 22 1271 Mössbauer spectra of NaI0 4 and Na 3 H2lOe; P. Jung and W. Triftshäuser, Phys. Rev. 175 (1968) 512. 23 i29i Mössbauer spectrum of K I 0 4 ; D. W. Hafemeister, G. de Pasquali and H. de Waard, Phys. Rev. 135B (1964) 1089. 24 Spectrum of NaC10 4 ; A. Fahlman, R. Carlsson and K. Siegbahn, Arkiv. Kemi, 25(1966)301. 25 Spectrum of K I 0 4 ; C. S. Fadley, S. B. M. Hagstrom, M. P. Klein and D. A. Shirley, / . Chem. Phys. 48 (1968) 3779. 26 Spectrum of NaC10 4 ; V.I.Nefedov,/.S/raci. CAem. 8 (1967) 919; W.Nefedow, Phys.Stat. Solidi, 2 (1962) 904. 27 Spectrum of NaC10 4 ; J. A. Bear den, Rev. Mod. Phys. 39 (1961) 86; D.S.Urch, / . Chem. Soc. (A) (1969) 3026. 28 Decomposition of KC10 4 ; W. H. Johnson and A. A. Gilliland, / . Res. Nat. Bur. Stand. 65A (1961) 63. 29 Decomposition of KBr0 4 ; G. K. Johnson, P. N. Smith, E. H. Appelman and W. N. Hubbard, Inorg. Chem. 9 (1970) 119. 30 Decomposition of alkali-metal perchlorates; M. M. Markowitz and D . A. Boryta, / . Phys. Chem. 69 (1965) 1114. 31 Decomposition of perchlorates of Mg, Ba, Cd, Zn and Pb; F. Solymosi, Magy. Kern. Foly. 74 (1968) 155. 32 Decomposition of perchlorates and metaperiodates of K, Rb and Cs; O..N. Breusov, N. I. Kashina and T. V. Revzina, Russ. J. Inorg. Chem. 15 (1970) 316. 33 For a review of this subject see M. Drätovsky and L. Pacesovä, Russ. Chem. Rev. 38 (1968) 243.

constants, reflecting the decrease in crowding around the central atom, are much more marked (Table 75). The frequencies of the first fully allowed electronic transitions ^ ( Ί 5 ^ 1 ) <- ^ ι θ ι 6 ) increase in the order IO4- < Br0 4 ~ < CIO4-. Perbromates are thermodynamically less stable than comparable perchlorates and periodates, and perbromic acid is potentially as strong an oxidizing agent as periodic acid (Tables 75 and 80). Comparison of the electronic structures and stabilities of CIO4-, Br0 4 ~ and IO4 ~ using Hartree-Fock-Slater atomic wave functions and self-consistent charge and configuration (extended Hückel) MO calculations568 shows that, although the total overlap population in Br0 4 - is slightly less than that in CIO4 - or IO4 ~, none of the earlier ideas about the "non-existence" of perbromates (high s-p promotion energy, poor d-p overlap or high inner-shell repulsions) was well founded; neither is extra stability conferred on the Κ>4~ ion by coulombic interactions or 4/-orbital participation. Inclusion of (n + l)s and nd orbitals gives charges for the halogen atoms of +0-57 for Cl, +0-65 for Br and +0-75 for I; for CIO4 ~ this is somewhat lower than the value obtained from spectroscopic techniques (Table 44), while investigations of IO4- (Table 76) have yielded charges for the iodine atom of +1 -24 (E.S.C.A.) and +1 -44 (Mössbauer effect).

1441

OXYACIDS AND OXYSALTS OF THE HALOGENS

Perchloric Acid and the Perchlorates 1. Preparation. The production of perchlorates as primary products may be accom­ plished in a number of ways734: (1) Electrolytic oxidation of chloride or chlorate to perchlorate in aqueous solution. (2) Action of a strong mineral acid (e.g. H 2 S0 4 ) on a chlorate, to give perchlorate, chloride and chlorine dioxide. (3) Thermal disproportionation of an alkali-metal chlorate, e.g. KCIO3 at ca. 500°C. (4) Chemical oxidation of chlorates to perchlorates with, for example, ozone, peroxydisulphates or lead dioxide. While the chemical methods are useful in the laboratory, only the electrolytic oxidation of chlorate is important commercially. The cells have a steel cathode and platinum anode, and the pH of the electrolyte is maintained at ca. 6-7; losses by reduction are minimized by the addition of chromate, which supposedly forms an insoluble film on the cathode. The predominant practice is to make sodium perchlorate first and thence other perchlorates by metathesis (Scheme 12). Ntfaor

Mg(C104)2,6H20

*- CCijOClOj

NaCIO,

FXeOClO.

AgClO^Hj

Fe(Cl04)x x=2 or 3

Se(OH)3C104

NO+ClO~

HC104,2S03

SCHEME 12. Interconversions of perchlorates.

Because of the size and limited coordinating power of the anion, many metal perchlor­ ates crystallize from solution with hydrated cations; the synthesis of anhydrous perchlor­ ates has been attempted by several routes. The low solubility of metal salts in anhydrous HCIO4 restricts one possible method. Some success has been secured in the reaction of NO +C104 ~ with metal salts: MX«+wNO + CIO4- -► M(C10 4 )«+/iNOX

(X = F, Cl, Br or NO3)

The high solubility of AgC104 in organic solvents has been utilized in synthesizing per­ chlorates of non-metallic cations, and also in preparing esters of perchloric acid.

1442

CHLORINE, BROMINE, IODINE AND ASTATINE: A. J. DOWNS AND C. J. ADAMS

2. Perchloric Acid Preparation1^ Aqueous solutions of perchloric acid are obtained by treating anhydrous NaC104 or Ba(C104)2 with concentrated hydrochloric acid, filtering off the precipitated metal chloride, and concentrating the filtrate by distillation. Depending on the reaction conditions, dehy­ dration of the constant boiling acid (72-4% HC104) affords either C1207 (p. 1374) or anhy­ drous HC104 as the major product; contamination of HCIO4 with C1207 may be minimized, but not eliminated, by using a modest ratio (ca. 3:1) of dehydrating agent (oleum or Mg(C104)2) to acid, and by immediately distilling the mixture at room temperature and very low pressure. Pure HCIO4 is best prepared by distilling the monohydrate H3O +C104 ~, obtained by treating crude anhydrous HC104 with the stoichiometric amount of the 72% acid (not with water). Deuteroperchloric acid has been obtained by distilling mixtures of D 2 S0 4 with NaC104 or Mg(C104)2. Properties Some properties of anhydrous perchloric acid are listed in Table 77; more extensive tabulations will be found elsewhere734»736»737. The stability of the acid is adversely affected by impurities, especially C1207, and many early studies of its chemistry are unreliable because the presence of C1207 went unrecognized. Molecular HOCIO3 has been defined in the vapour phase by electron diffraction (Table 77); infrared and Raman spectroscopic results affirm that the anhydrous acid retains this molecular form in the solid and liquid phases, although there is evidence of association via O-H · · · O hydrogen bonds. Recent investigations of some physical properties of liquid HC104 (vapour pressure, viscosity, electrical conductivity and electrolysis)741*742 suggest that the dissociation 3HC104 # C 1 2 0 7 + H 3 0 + + C K V

occurs with an equilibrium constant at 25°C of 0-68 x 10 ~6. Anhydrous perchloric acid is an extremely powerful oxidizing agent; it reacts explosively with most organic substances, ignites hydrogen iodide and thionyl chloride, and rapidly oxidizes gold and silver. Some metal salts dissolve with formation of the anhydrous metal perphlorate, e.g. A12C16+6HCIO4 -> 2A1(C104)3+6HC1 Mn(N0 3 ) 2 +6HC10 4 -► Mn(C10 4 )2+2N02 + C10 4 - + 2 H 3 0 + C 1 0 4 -

Solutions of HC104 in non-oxidizable organic liquids (e.g. CHC13, MeCN or MeN0 2 ) are best obtained by passing dry HC1 gas into the appropriate solution of silver perchlorate. In contract with the hypohalous acids, HOCIO3 does not add across olefinic double bonds. Aqueous Solutions Perchloric acid forms a constant boiling mixture with water which contains 72-4% by weight of the acid and boils at 203°C with slight decomposition. At room temperature aqueous solutions have the properties only of a very strong acid, dissolving active metals with liberation of hydrogen and forming substituted ammonium salts with amines. 79% perchloric acid has an indicator activity equivalent to that of 98% sulphuric acid. Spectro­ scopic studies (Table 78) show dissociation of the acid to be complete at concentrations < 6 M, ™ V. Ya. Rosolovskii, Russ. J. Inorg. Chem. 11 (1966) 1158. 742 N. Bout and J. Potier, Rev. Chim. Minerale, 4 (1967) 621.

OXYACIDS AND OXYSALTS OF THE HALOGENS

1443

TABLE 77. PROPERTIES OF ANHYDROUS PERCHLORIC ACID

Thermodynamic and other physical properties Δ # / ° for HClO 4 (0 at 25°C - 9 - 7 0 kcal ΓηοΓ 1 a S° for HC10 4 (g) at 25°C 68-2caldeg-imol-lb Cp° for HCIO4O) at 25°C 28-8caldeg-imorib Colour colourless when freshly distilled Melting point -101°C*> 1-657 kcal m o l - i b A/Zfuslon 120-5°CC Boiling point 813kcalmol-ic ΔΗνΛΡ at the boiling point 20-6caldeg-imol-ic Trouton's constant Ref. c Vapour pressure 1-761 g c m - 3 d Density of the liquid at 25°C Ref. e Electrical conductivity Molecular structure Measured at 35°Cf Electron diffraction: 1-408(2) A r,(l)[Cl = 0 ] 1-635(7)A r,(D[Cl-OH] 112-8(5)° 105-8(7)° L 0 = Cl~OH Spectroscopic properties Measured for HCIO4 and DCIO4 in the Infrared spectrum solid, liquid and vapour phases* Measured for the liquid h and solid1 Raman spectrum Vibrational assignment of skeletal fundamentals v8ymm(C103), 1036; KC1—OH), 742; 8symm(C103), 577; vaeymm (CK> 3), 1230; Saeymm(C103), 577; P(C\03), 430 cm" i 9-20mdyneA-i Force constants: / r (Cl = 0 ) 3-85mdyneA-i Λ(α-ΟΗ) Refs.j,k, 1 Mass spectrum 13-05+0-1 eV, measured by electron Ionization potential impact k Bond dissociation energy: 46kcalmol~11 Z)(HO—CIO3)

z.o=ci=o

a Selected Values of Chemical Thermodynamic Properties, N.B.S. Technical Note 270-3, Washington (1968). b J. C. Trowbridge and E. F . Westrum, jun., / . Phys. Chem. 68 (1964) 42. c D . J. Sibbett and I. Geller, cited in G. S. Pearson, Ado, Inorg. Chem. Radiochem. 8 (1966) 177. d G. Mascherpa, Bull. Soc. chim. France (1961) 1259. * N . Bout and J. Potier, Rev. Chim. Mintrale, 4 (1967) 621. f A. H. Clark, B. Beagley, D . W. J. Cruickshank and T. G. Hewitt, / . Chem. Soc. (A) (1970) 1613. « P. A. Giguere and R. Savoie, Canad. J. Chem. 40 (1962) 495. h A. Simon and M. Weist, Z . anorg. Chem. 263 (1952) 301. 1 A. J. Dahl, J. C. Trowbridge and R. C. Taylor, Inorg. Chem. 2 (1963) 654. i H. F . Cordes and S. R. Smith, / . Chem. Eng. Data, 15 (1970) 158. k I. P. Fisher, Trans. Faraday Soc. 63 (1967) 684. 1 G. A. Heath and J. R. Majer, Trans. Faraday Soc. 60 (1964) 1783.

and not less than 90% at 12 M. In 6 M solution there are four water molecules for every ionic species in solution, and incomplete ionization at higher concentration is a result of the limited number of water molecules available to solvate the proton. In accord with these obser­ vations, perchloric acid is normally extracted from water into ether as [(^O^m+fClOJ-".

1444

CHLORINE, BROMINE, IODINE AND ASTATINE: A. J. DOWNS AND C. J. ADAMS

Perchloric acid solutions exhibit very little oxidizing power in the cold; they are slowly reduced by SnCl2, V 2 0 3 and V(II), V(III) and Ti(III) ions. Rates of reaction with V 2 0 3 fall in the sequence I 0 4 _ > Br0 3 ~ > I0 3 ~ > CIO4-. The hot concentrated acid is a vigorous oxidizing agent, however, evolving hydrogen chloride with metals and decomposing organic materials; the "wet fire" method (used for ashing organic material prior to the determination of inorganic residues) utilizes perchloric acid alone or in mixtures with nitric or periodic acid. TABLE 78. PROPERTIES OF AQUEOUS PERCHLORIC ACID

Thermodynamic properties at 298°K MIf° AG/° S° A/f d i lu tion(HC104->HC104,ooH 2 0) Electrical conductivity Acidity Spectroscopic studies of dissociation for [HCIO4] < 12 M Hammett acidity function, H0

- 30-91 kcal mol" 1 a - 2 0 6 kcal m o l - i » 43-5 cal deg"i mol" 1 a -21-21 kcalmol-i · Ref. b Raman studies c - d H nmr studies c'e»f 35 C1 nmr studies c Experiment « Calculated h 1

a Selected Values of Chemical Thermodynamic Properties, N.B.S. Technical Note 270-3, Washington (1968). to L. H. Brickwedde, / . Res. Nat. Bur. Stand. 42A (1949) 309. c J. W. Akitt, A. K. Covington, J. G. Freeman and T. H. Lilley, Trans. Faraday Soc. 65 (1969) 2701. d K. Heinzinger and R. E. Weston, jun.,/. Chem.Phys. 42 (1965) 272. e R. W. Duerst, / . Chem. Phys. 48 (1968) 2275. *G. C. Hood and C. A. Reilly, / . Chem. Phys. 32 (1960) 127. * Measured using substituted aniline indicators; K. Yates and H. Wai, / . Amer. Chem. Soc. 86 (1964) 5408. h J. G. Dawber, Chem. Comm. (1966) 3.

Hydrates of Perchloric Acid The melting point diagram of the system CI2O7-H2O indicates the existence of at least nine intermediate phases743, several of which have been characterized by diffractometric and spectroscopic techniques (Table 79) as containing C10 4 _ ions hydrogen-bonded to [H(OH2)x]+ cations. The monohydrate H 3 0 + C104~ has been particularly well studied. The phase transition at 243°K is associated with a change in the rotation of the H3O + ion in the lattice; proton spin-relaxation times for the polycrystalline solid demonstrate that below 243°K the H3O + ion is rotating about a threefold axis, but that above 243°K the reorienta­ tion is essentially isotropic. The CIO4 ~ anion in the orthorhombic high-temperature form is essentially tetrahedral, but the monoclinic low-temperature modification contains slightly distorted C104~ anions such as have been found in the other hydrates; Cl-O bonds longer than the mean are involved in hydrogen-bonding. The H 5 0 2 + cation in HC104,2H20 is centrosymmetric. Decomposition of Anhydrous Perchloric Acid In contrast with its hydrates and aqueous solutions—72% HCIO4 boils at 203°C with 743 G. Mascherpa, Rev. Chim. Minirale, 2 (1965) 379.

—

-100·

Decomposes at -73-1·

(HC104)4,H20

ca. -30* -91-35· at 25°Cf»* at -80°C* at temperatures in the range -50° to 25°0 at -180οΟ> and 25°C at temperatures in the range-180° to 25°Cd at temperatures in therange-180° to 25°0» at - lS0°O*

at - Ι δ Ο ^

-162-04« SLt-196°0

—

at -188°C J

b c d

—

at-188°C*

[H703]+[C104]-40-2±0-l°

[H 3 0] + 2 [C10 4 ]- 2 ,3H 2 0 -33-l±0-l c

[H502]+[C10c4]-20-65±0-l

[H30]+[C104]+49-905* 203

HC104,3H20

HC104,5/2H20

HC104,2H20

HC104,H20

Composition

• G. Mascherpa, Compt. rend. 252 (1961) 1800. G. F. Smith and O. E. Goehler, Ind. Eng. Chem. Anal. Edn. 3 (1931) 61. G. Mascherpa, A. C. Pavia and A. Potier, Compt. rend. 254 (1962) 3215. D. E. O'Reilly, E. M. Peterson and J. M. Williams, / . Chem. Phys. 54 (1971) 96. • Selected Values of Chemical Thermodynamic Properties, N.B.S. Technical Note 270-3, Washington (1968). 1 Average KCl-O) = 1 -452(5) A; M. R. Truter, Acta Cryst. 14 (1961) 318. • Average KCl-O) = 1 -42(1) A; F. S. Lee and G. B. Carpenter, / . Phys. Chem. 63 (1959) 279. • Average KCl-O) = 1 -46i(6) A; C. E. Nordman, Acta Cryst. 15 (1962) 18. 1 Average KCl-O) = 1-438(3) A; I. Olovsson, / . Chem. Phys. 49 (1968) 1063. 1 Average KCl-O) = 1 -434(2) A; 1 -439(2) A; J. Ahmlöf, J.-O. Lundgren and I. Olovsson, Acta Cryst. B27 (1971) 898. k Average KCl-O) = 1-437(2) A; J. Ahmlöf, Acta Cryst. B28 (1972) 481. 1 R. C. Taylor and G. L. Vidale, / . Amer. Chem. Soc. 78 (1956) 5999. m A. C. Pavia and P. A. Giguere, / . Chem. Phys. 52 (1970) 3551. • R. Savoie and P. A. Giguere, / . Chem. Phys. 41 (1964) 2698.

35C1 nqr spectrum

*H nmr spectrum

Infrared spectrum

Raman spectrum

Boiling point (°C) Transition temperature (ifany)(°C) AÄ / °at25°C(kcalmol-i) X-ray diffraction

Structure Melting point (°Q

Property

TABLE 79. PROPERTIES OF THE HYDRATES OF PERCHLORIC ACID

—

—

-45-9±0-l*

HC104,7/2H20

1446

CHLORINE, BROMINE, IODINE AND ASTATINE: A. J. DOWNS AND C. J. ADAMS

only slight decomposition—anhydrous HC104 is apparently unstable even at — 78°C744. The explosive reputation of HCIO4 itself is probably exaggerated, at least for very pure samples; however, the acid ages badly, and the presence of impurities, especially C1207, greatly facilitates decomposition. Pyrolysis of perchloric acid vapour 4HC10 4 -+ 2 H 2 0 + 2 C 1 2 + 7 0 2

has been studied in the temperature range 15(M39°C745»746. Above 310°C the reaction is homogeneous and first-order in HC104; the suggested rate-determining step is rupture of the Cl-OH bond, HOC10 3 ->HO+C10 3

followed by a fast reaction between a hydroxyl radical and another perchloric acid molecule HO 4- HOCIO3 -► H 2 0+CIO4

and the decomposition of CIO3 and CIO4 to chlorine and oxygen via C102 and ClO. The experimental activation energy (45-1 kcal mol -1 ) accords well with the dissociation energy of the Cl-OH bond measured by electron impact (46 kcal mol -1)747 0 r calculated from thermodynamic data (47-6 kcal mol -1 ). Below 310°C the decomposition is heterogeneous and second-order in HCIO4, while at temperatures well above 450°C hydrogen chloride is produced by reaction between chlorine and water: C1 2 +H 2 0->2HC1+K>2

On standing at room temperature, the initially colourless, mobile, hygroscopic liquid discolours, passing through yellow and red to become dark brown in a matter of 2 days; at this point evolution of oxygen and chlorine oxides sets in, and after 4 days the liquid, now colourless again, deposits white crystals of H3O +CIO4 ~. Again the initial step is cleavage of the Cl-OH bond. The instability of anhydrous HCIO4 compared with its solutions is caused by the relative weakness of the Cl-OH bond, which is of course absent in the dissociated acid. The presence of only small amounts of dichlorine heptoxide significantly alters the character of the decomposition. Other Systems Involving Perchloric Acid Perchloric acid is sufficiently strong an acid to protonate many molecules which are themselves acidic in aqueous solution: selenous acid and perchloric acid form the compound [Se(OH)3]+C104-, while mixtures of anhydrous nitric and perchloric acids at -40°C give evidence for the equilibrium *

HNO3+2HCIO4 ^ NCVCKXr +H30+C104-

Physical properties have been reported for systems involving perchloric acid and the follow­ ing acids734,735,736: CH„Cl3_wC02H (n = 0, 1, 2 or 3), H 2 S0 4 , ΗΝ0 2 , Η3ΡΟ4 and CF3C02H748. 744 A. A. Zinov'ev, Zhur. Neorg. Khim. 3 (1958) 1205. 745 j . B . Levy, / . Phys. Chem. 66 (1962) 1092. 746 D . J. Sibbett and I. Geller, cited in G. S. Pearson, Adv. Inorg. Chem. Radiochem. 8 (1966) 177. 747 G. A. Heath and J. R. Majer, Trans. Faraday Soc. 60 (1964) 1783.

1447 Perchloric acid is apparently completely dissociated in solvents such as acetonitrile. In organic acids, however, it is only moderately strong; for example, the pKa value is ca. 0-5 in formic acid, 2-7 in acetic acid and ca. 1 -3 in trifluoroacetic acid. Perchloric acid has been used in acetic acid (or acetic acid-acetic anhydride mixtures) for the potentiometric or conductimetric titration of organic bases749»750, such as amines or phosphine oxides, which give only diffuse end-points in aqueous solution. The method has been extended to the analysis of ketones by titration as semicarbazone or hydrazone derivatives751, and other solvents have been used, including chlorobenzene and acetonitrile. Infrared and Raman investigations of solutions of HC104 in acetic anhydride have shown752»753 that the equilibrium OXYACIDS AND OXYSALTS OF THE HALOGENS

(CH 3 CO) 2 0+HOC10 3 ^ CH 3 C(0)OC103+CH3C0 2 H

is established. Acetyl perchlorate has also been prepared from silver perchlorate and acetyl chloride; it is a vigorous acetylating agent. The benzoyl analogue is also known. 3, Ionic Perchlorates. The many perchlorate salts which have been isolated contain both inorganic and organic cations. The latter compounds, although unstable thermodynamically, are isolable at room temperature because of the kinetic inertness of the perchlorate ion; they may be dangerously explosive if subjected to shock or heating. Perchloric acid forms salts BH + C10 4 - by protonating weak organic bases B (amines, amine oxides or ketones); carbonium and diazonium perchlorates are also known. Non-metallic inorganic cations include NO+, N0 2 + and N2H5+. The large size of the perchlorate ion regulates many of the properties and uses of the compounds. Thus, many metal perchlorates crystallize with hydrated cations, and ammines of perchlorates have been reported. Anhydrous perchlorates, e.g. Mg(QC>4)2, are frequently deliquescent, and absorb not only water but many other polar molecules from the vapour phase. Perchlorates of large unipositive cations tend to be insoluble in water; that KCIO4 is far less soluble than NaC104 has been utilized in the separation of these alkali metals, while complex cations are often conveniently crystallized as perchlorates for subsequent investi­ gation by X-ray diffraction. By contrast, the perchlorates of lithium, sodium, the alkaline earths and especially silver are extremely soluble in organic media. Silver perchlorate is used in many metathetical preparations of other perchlorates (Scheme 12), and forms crys­ talline complexes AgC104,L, AgC104,2L and 4AgC104,L (L = aromatic hydrocarbon) in which the organic molecule is bonded to the cation; some of these compounds have been studied crystallographically, e.g. AgC104,C6H6 75*, AgC104,2(m-xylene)755 and 4AgC104,L (L = naphthalene or anthracene)756. The relation between the relative sizes of counter-ions, stoichiometry and the solubility of salts has been discussed in terms of the variation in lattice and solvation energies757. 748 j . Bessere, Bull. Soc. chim. France, (1969) 3356. 749 A . H. Beckett and E. H. Tinley, Titration in Non-aqueous Solvents, 3rd edn., B.D.H., Poole (1960); I. M. Kolthoif and S. Bruckenstein, Treatise on Analytical Chemistry (ed. I. M. Kolthoff, P. J. Elving and E. B. Sandell), Part I, Vol. I, p. 475, Interscience, New York (1959). 750 v . Vajgand and T. Pastor, / . Electroanal. Chem. 8 (1964) 40. 751 D . B. Cowell and B. D . Selby, Analyst, 88 (1963) 974. 752 A.-M. Avedikian and A. Commeyras, Bull. Soc. chim. France, (1970) 1258. 753 E . S. Sorokin, N. I. Geidel'man and V. Ya. Bytenskii, Zhur. priklad. Khim. 43 (1970) 1595. 754 H . G. Smith and R. E. Rundle, / . Amer. Chem. Soc. 80 (1958) 5075. 7551. F . Taylor, jun., E. A. Hall and E. L. Amma, / . Amer. Chem. Soc. 91 (1969) 5745. 756 E . A. Hall and E. L. Amma, / . Amer. Chem. Soc. 91 (1969) 6538. 757 D. A. Johnson, Some Thermodynamic Aspects of Inorganic Chemistry, pp. 97-115, Cambridge (1968).

1448

CHLORINE, BROMINE, IODINE AND ASTATINE: A. J. DOWNS AND C. J. ADAMS

Structural Aspects Essentially tetrahedral C104" anions have been defined crystallographically in the salts listed in Table 76, though rather more precise dimensions have been obtained for the crystalline hydrates of perchloric acid (Table 79). The bond lengths, uncorrected for thermal motion, are ca. 1-44 A; including the correction increases the bond lengths to ca. 1-46 A. Many early studies of the perchlorates were inaccurate, and these have not been listed; likewise excluded from the list are some more recent crystallographic investigations of perchlorates containing complex cations, which achieved their main aim of characterizing the cation at the expense of reporting chemically unrealistic parameters for the CIO4- unit. Coordination of perchlorate groups to metal ions has been detected by vibrational spectroscopy: while the Raman spectra of concentrated aqueous solutions show no effects not attributable to ion-pairing758, the infrared spectra of some complex transition metal perchlorates evince coordination of the CIO4 ~ unit to the cation. The discrete tetrahedral C104~ ion has four characteristic Raman lines (Table 75), two of which coincide with dis­ tinctive infrared absorptions at ca. 1050-1150 c m - 1 (v3) and 630 c m - 1 (v4); the totally symmetric stretching fundamental, although formally forbidden, is often observed as a weak band at ca. 930 cm - 1 in the infrared spectra of the solids. If the perchlorate ion becomes coordinated to a metal ion by one oxygen atom, the local symmetry of the — OCIO3 group is reduced to C$v (with the assumption of a linear M-O-Cl unit). As a result, the broad degenerate v3 band splits into two well-defined features between 1000 and 1200 cm _1 ; stretching of the coordinated Cl-0 bond (formally related to vx of CIO4-) is seen as an intense infrared absorption at 925-950 cm - 1 , while the degenerate bending mode v4 also splits into two bands. Such coordination has been found for perchlorates of some divalent metals, including /rawj-Co[o-phenylenebisdimethylarsine]2(C104)2 759, Ni(MeCN)x(C104)2 (x = 2 or 4) 760 and NiL2(C104)2 (where L = N-methylethylenediamine or N,N'-dimethylethylenediamine)761. The magnitude of the band splittings observed is apparently consistent with a force constant of ca. 6 mdyne A - 1 for the coordinated Cl-O bond, as opposed to ca. 8 mdyne A - 1 for the terminal Cl-O bonds762. Bidentate perchlorate groups, characterized by further changes in the infrared spectrum, have been suggested for the compounds NiL2(C104)2 (L = Ν,Ν,Ν'-trimethylethylenediamine)76i, Ni(ethylenediamine)2(C104)2 763 and U0 2 (C10 4 ) 2 764. Thermal Decomposition Depending on the cation, metal perchlorates decompose to yield the corresponding oxide or chloride; in most cases the reaction achieves a measurable rate below 600°C. M(C10 4 )n->MCln+2/i0 2 M(C104)n -► MO« /2 +/i/2Cl2+7/i/402

The manner of decomposition may be rationalized by considering the relative equivalent free energies of the possible products765. The perchlorates of the alkali metals, silver, calcium, barium, cadmium and lead decompose to afford the corresponding chlorides; for 758 R. E . Hester and R. A . Plane, Inorg. Chem. 3 (1964) 769. 759 G . A . Rodley and P. W . Smith, / . Chem. Soc. (A) (1967) 1580. 760 A . E . Wickenden and R . A . Krause, Inorg. Chem. 4 (1965) 4 0 4 . 761 S. F . Pavkovic and D . W . Meek, Inorg. Chem. 4 (1965) 1091. 762 H . Brintzinger and R. E . Hester, Inorg. Chem. 5 (1966) 980. 763 M . E . Farago, J. M . James and V. C . G . Trew, / . Chem. Soc. (A) (1967) 820. 764 v . M . Vdovenko, L . G . Mashirov and D . N . Suglobov, Soviet Radiochemistry, 9 (1967) 37. 765 M . M . Markowitz, / . Inorg. Nuclear Chem. 2 5 (1963) 407.

OXYACIDS AND OXYSALTS OF THE HALOGENS

1449

-1

these elements, AGy°(chloride) - JAGy°(oxide) < - 2 0 kcal mol . Aluminium and ferric perchlorates, on the other hand, produce the corresponding oxides (here AGy°(chloride) — JAGy°(oxide) > 0 kcal mol -1 ), while for magnesium and zinc both reactions are observed. Dehydration of hydrated perchlorates may be difficult, commonly being accompanied by decomposition, in the event of which hydrolysis reactions may also occur. Decomposition of the alkali-metal perchlorates involves little heat change: MC10 4 -> MC1+20 2

AH° ~ 0

While the corresponding chlorates have long liquid ranges, rapid decomposition of the heavier alkali-metal perchlorates begins at ca. 580°C and apparently causes melting; the exact temperature depends on particle size and on the previous history of the sample. Only LiC104 has a congruent melting point (at 247°C). There is no appreciable accumulation of intermediate chlorates, which are, of course, highly unstable at the temperatures involved; by contrast, formation of the thermally more stable halate is a well-defined step in the pyrolysis of alkali-metal perbromates and metaperiodates (Table 73). The activation energy for perchlorate decomposition (corresponding to rupture of a Cl-O bond) decreases from 60-70 kcal mol - 1 for the alkali-metal perchlorates to 40 kcal mol - 1 for Zn(C104)2. The thermal decomposition of ammonium perchlorate has been extensively studied, not only for its chemical interest, but also because of its application in the oxidation of rocket fuels; there have been three recent reviews766»767»768. The decomposition was first observed in 1869 when the simple reaction NH4CIO4 ->NH 4 Cl+20 2 was proposed; it has since been realized, however, that three fairly distinct stages are involved: (1) "Low-temperature" decomposition in the range ca. 200-300°C, which is characterized by an induction period, an acceleratory region, a rate maximum and a deceleratory region, the reaction stopping before all the material is consumed: the approximate stoichiometry is 4NH4CIO4 -> 2CI2+8H2O+2N2O+3O2

(2) "High-temperature" decomposition, between 350°C and 400°C, with immeasurably fast initiation, the stoichiometry being 2NH4CIO4 - * α 2 + 4 Η 2 0 + 2 Ν Ο + 0 2

(3) Deflagration at ca. 450°C, for which two limiting equations have been found: (a) at low pressure: 2NH4CIO4 - * C l 2 + 4 H 2 0 + 2 N O + 0 2

(b) at high pressure: 4NH4CIO4 -> 4HC1+ 6 H 2 0 + 2 N 2 + 5 0 2

Other minor products reported in individual studies include N 2 0, N2O3, NO2, N 2 0 4 , N 2 , HC1, C102, NOC1, N02C1, HNO3 and HC104. Careful investigation has shown that the reaction isfirst-orderand that the activation energy {ca. 30 kcal mol- 1 ) is the same in all three regions; the rate-determining step is believed to be proton-transfer from NH 4 + to C104 - in the crystal, followed by evaporation into the gas phase. Decomposition of HC104 and subsequent oxidation of NH 3 by the fragmentation products in vapour-phase reactions account for the majority of the observed 766 A. G. Keenan and R. F. Siegmund, Quart. Rev. Chem. Soc. 23 (1969) 430. 767 v . V. Boldyrev, Dokl. Akad. Nauk S.S.S.R. 181 (1968) 1406. 768 p. w . M. Jacobs and H. M. Whitehead, Chem. Rev. 69 (1969) 551.

1450

CHLORINE, BROMINE, IODINE AND ASTATINE: A. J. DOWNS AND C. J. ADAMS

products, whose distribution depends on the reaction conditions. NH^C10;(s)

«

— NH 3 (ads.)

sublimate

~«

NH 3 (g)

tl

+

+

HC104(ads.) —

m* products

HC104(g)

+~ products

II

Radiolytic Decomposition The action of X-rays and y-rays on the alkali-metal and alkaline earth perchlorates has been comprehensively studied769. Chemically identifiable fragments include the metal oxide (or Superoxide), C103 -, C102, C102 -, ClO -, Cl - and 0 2 ; 0 3 ~, ClOO, C103 and C104 have been detected by esr spectroscopy as paramagnetic centres stable at low tempera­ ture610. Similar results have been obtained in the radiolysis of frozen or liquid aqueous solutions of perchloric acid, although these studies are complicated by the reaction of radical intermediates with water736. The mechanism is complex, but probably involves several of the possible primary decomposition processes as well as excitation of C104~.

It is interesting to note that ClOO is a. precursor of chlorine dioxide, and that C104 probably has the structure 0 2 C100 6 1 0 ; Cl ~ is a primary product in radiolysis of the solid, whereas in solution it arises from the attack on C103 ~ by various free radicals. 4. Esters of Perchloric Acid Alkyl Perchlorates Ever since their discovery in 1841 the alkyl perchlorates have been recognized as treacherously explosive compounds; their sensitivity to shock decreases as the molecular weight rises, although the higher temperatures needed to distil the compounds of higher molecular weight introduce a compensatory hazard. The perchlorate esters have been stabilized in the form of urea inclusion compounds770. Although the oily liquids may be obtained from the anhydrous acid and the alcohol at low temperature, the best preparative procedure involves metathesis between silver per­ chlorate and an alkyl halide in an organic solvent. The infrared spectra of the liquids are consistent with the structure ROCIO3 7 7 °; features at ca. 705, 1035, 1230 and 1260 cm-i are attributed to vibrations of the — OCIO3 group. Studies have been made of the substitution reactions of methyl perchlorate. Hydrolysis involves attack at both the chlorine and carbon atoms, while methanolysis is catalysed by C10 4 - 7 7 i: MeOClOs+MeOH - * Me 2 0+HOCIO3 769 L . A . Prince and E . R. Johnson, / . Phys. Chem. 6 9 (1965) 359, 377. 770 j . Radell, J. W . Connolly and A . J. R a y m o n d , / . Amer. Chem. Soc. 8 3 (1961) 3958. 771D. N . Kevill and H . S. Posselt, Chem. Comm. (1967) 4 3 8 ; J. Koskikallio, Suomen Kemistilehti, (1967) 199.

40B

OXYACIDS AND OXYSALTS OF THE HALOGENS

1451

Extremely explosive materials have been made by partial or complete reaction of anhydrous perchloric acid with polyfunctional alcohols including glycol, glycerol and pentaerythritol; explosions have occurred merely on pouring the liquids from one container to another (cf. nitroglycerine). The compound CCI3OCIO3 explodes on contact with organic material. Inorganic Molecular Perchlorates Inorganic analogues of the esters of perchloric acid include the halogen perchlorates FOCIO3, CIOCIO3, BrOC103 and I(OC103)3, which are discussed fully in Section 4B5, and the xenon compounds FXeOQC>3 and Xe(OC103)2, prepared from xenon difluoride and anhydrous perchloric acid601, and involving linear coordination about the xenon atom. 5. Uses In Explosives and Propellants1**'16* KCIO4 and NH4CIO4 are important oxygen-carriers for solid explosives and propellants; more powerful explosives contain ammonium perchlorate, since with suitably combustible material this gives entirely gaseous products: NH4CIO4 + 2C

induced explosion > NH3 + HCI + 2CO2

Perchlorate explosives are not sensitive to shock and are capable of wide modification to suit the particular purpose in hand. Perchlorates have also been included in slower-burning mixtures used for fireworks or signal flares; heavy metal salts are used to introduce colours. Typical rocket propellants, whose combustion might be described as a slow explosion, consist of 75% NH4CIO4, 20% fuel (which also serves to bind the material into pellets) and 5% additives to provide desired effects on physical properties, storage or burning capabili­ ties. Lithium perchlorate has recently been used as an oxidizer since it has twice as much oxygen per unit volume as NH4CIO4 and gives a higher combustion temperature. Other Uses Several analytical uses of perchloric acid—as an acid in non-aqueous titrimetry and as an oxidant in the wet ashing of organic material—have been mentioned in the preceding text. Perchloric acid has also been used for the digestion of chromium, including its quanti­ tative separation from other metals, as an electrolyte (together with acetic acid) in the electro-polishing of aluminium, and as a catalyst for esterification reactions. Miscellaneous uses are listed elsewhere736. Perchlorates have been used as drying agents: anhydrous Mg(C104)2 can absorb up to 60% of its own weight of water, and does not become sticky upon handling; regeneration can be achieved in vacuo at 200°C. The minimal complexing ability of the CIO4- anion in aqueous solution has led to the widespread adoption of perchlorates to provide media of constant ionic strength for the study of the kinetics and equilibrium constants of many aqueous processes. Perbromic Acid and the Perbromates575b·663 First obtained in studies of the /?-decay of 83 Se04 2_ 663, perbromates are also formed in the y-radiolysis of crystalline bromates725. Oxidation of bromate ions in aqueous solution is

1452

CHLORINE, BROMINE, IODINE AND ASTATINE: A. J. DOWNS AND C. J. ADAMS

effected electrolytically663, with XeF 2 663, or, most expeditiously, by passing fluorine into strongly basic solutions containing bromate731. Perbromic acid is a strong acid bearing one proton per heptavalent bromine atom; in aqueous solution it is completely dissociated into aquated protons and tetrahedral BrC>4 ~ anions, which do not exchange oxygen with H 2 1 8 0. Solutions more than 6 M (55%) in HBrC>4 are unstable in air, undergoing an apparently autocatalytic decomposition which is complete for 80% concentrations: decomposition is also catalysed by metal ions, e.g. Ag + and Ce4 + . Despite this decomposition, some perbromic acid may be distilled, and the mass spectrum of HBrC>4 has been reported772. On very rapid evaporation, crystallization of perbromic acid solutions occurs (possibly to give HBr04,2H 2 0) just before decomposition sets in. Attempts to dehydrate the acid have met only with decomposition. The vibrational spectra reported for the tetrahedral Br0 4 - anion vary but little, irrespec­ tive of whether the ion is in solution or in various crystalline salts. KBrC>4 crystallizes with the orthorhombic Barite structure773, being isomorphous with the room-temperature modifications of the alkali-metal perchlorates; the Br0 4 ~ unit is tetrahedral within the limits of experimental error, having a Br-O bond length of 1-61 Ä. KBr0 4 is converted by SbF 5 into Br0 3 F (p. 1393)653. The exothermic reaction KBr0 4 (c) -+ KBr(c)+20 2 (g)

Δ/Γ 2 9 8 = - 2 5 · 3 8 ± 0 · 1 0 kcal mol"i

proceeds in two stages: (i) production of KBr0 3 at ca. 275°C, and (ii) its subsequent decomposition at ca. 390°C774. From the calorimetric study of this decomposition, and of the dissolution of KBr04 in water, were derived the thermodynamic properties of KBr0 4 and Br0 4 ~(aq) listed in Tables 73 and 75™. The order of thermal stability KBr0 3 > KBr0 4 resembles that for the iodine systems rather than that for the analogous chlorine compounds (KC104 > KCIO3). Unlike NH4CIO4, NH 4 Br0 4 775 is neither shock- nor friction-sensitive; in its thermal decomposition (producing N 2 , Br2, 0 2 and H 2 0), it resembles NH 4 Br0 3 rather than NH4CIO4 (p. 1449). Despite the large potential ( + 1 -76 volts) of the Br0 4 ~/Br0 3 ~ couple in acidic solution, perbromic acid is described as a sluggish oxidizing agent at room temperature, though somewhat more powerful than perchloric acid: dilute solutions oxidize Br~ and I~ only slowly, but the 12 M acid rapidly oxidizes chloride. At 100°C 6 M perbromic acid fairly rapidly converts Mn 2 + to Mn0 4 ~. Reduction to bromide may be accomplished with SnCl2. Periodic Acid and the Periodates Nomenclature Despite efforts740 to establish a logical system for naming derivatives of H 5 IOÖ in various stages of deprotonation, dehydration and aggregation, the older nomenclature is still most commonly used, based on four acids: H 5 I0 6 , para- or ortho^trioaic acid; HIO4, //^aperiodic acid; (hypothetical) H 2 I0 5 , m&söperiodic acid; and H7I3Oi4, fr/periodic acid. Protonated anions are described thus: H 3 I0 6 2 _ , trihydrogen paraperiodate; but degrees of aggregation are difficult to render unambiguously using this system. In discussing periodates it is least confusing and cumbersome to use formulae whenever possible. 772 M . H . Studier, / . Amer. Chem. Soc. 9 0 (1968) 773 s . Siegel, B . Tani and E . H . Appelman, Inorg. 774 G . K. Johnson, P. N . Smith, E . H . Appelman 775 j . N . Keith and I. J. S o l o m o n , Inorg. Chem. 9

1901. Chem. 8 (1969) 1190. and W . N . Hubbard, Inorg. Chem. 9 (1970) 119. (1970) 1560.

1453

OXYACIDS AND OXYSALTS OF THE HALOGENS aq. KOH

KIO^ H,()

i

K4H,l:O|0.8H:O

70X

1

^ K

4

I

:

O

y

T<35°C NaJO z NalO,

dil. HNO, Lor NaI0 3

AgI0 4

cone. HNO,

Cl,

A Na

a q.NaOH

Nal + 2NaX>

35()°C

ΚΝΟ,,

3H2I06

gN°3L dil.HNÖ 3

60°C, ,ΠΓΗΝΟ,

Ag 2 H 3 I 0 6

Ag

acid ~

2HI°5

A l

h °>

KOH dil. Ba(NO) ? HNO,

Ag 5 I0 6 H l

7 3°,4

120°C> I0 2 F 3

(i)HSO.F , 3 -i Ba_[H,lO]

(ü)so3.

3

-

62

cone. HNO, 3

-Ba(N0 3 ) 2 t

H

5I06

100 712mm

HKX

SCHEME 13. Interconversion of periodates.

Synthesis The synthesis of periodates rests mainly on the oxidation of iodide, iodine or iodate in aqueous solution. Industrial processes employ electrochemical oxidation with a Pb0 2 anode. A starting point for the laboratory preparation of many periodates714 is the sodium salt Na3H2I06 (Scheme 13), conveniently produced by passing chlorine into aqueous alkaline solutions of iodine or sodium iodate714. The paraperiodates of the alkaline earth metals can be obtained by disproportionate of the iodates, e.g. 5Ba0O3)2 -> Ba5(I06)2+4I2+902

The thermal stability of paraperiodates also allows their synthesis by the oxidation of solid mixtures of iodides and oxides with oxygen. Periodic Acid Aqueous solutions of periodic acid are best obtained by treating barium paraperiodate with concentrated nitric acid: the only solid phase stable in contact with such a solution is paraperiodic acid, H 5 IOÖ (Table 80). Structural investigations of the white crystalline solid confirm that it is a genuine ortho acid (HO)5IO rather than a hydrate HI04,2H 2 0 or an oxonium salt [Η 5 θ2] + Ιθ4~ analogous to the dihydrate of perchloric acid (Table 79). Dehydration of H5IO6 on heating to ca. 120°C produces H7I3O14, while heating at 100°C in vacuo gives HIO4, which has also been characterized in the mass spectrum of periodic acid; at higher temperatures evolution of oxygen sets in, producing initially a solid I2O5J2O7 and, at 150°C, just I 2 0 5 . The reaction of H 5 I0 6 with 60% oleum is said to produce an orange solid of approximate composition I2O7 625. Periodic acid is a fairly weak acid, and in acid solution exists in the undissociated form (HO)5IO. The difference between the apparent (pK 1 -6) and true (pK 3-29)firstdissociation constants of H 5 IOÖ is caused by hydration equilibria between CO2 and the undissociated

1454

CHLORINE, BROMINE, IODINE AND ASTATINE: A. J. DOWNS AND C. J. ADAMS TABLE 80. PROPERTIES OF PERIODIC ACID, H 5 IOÖ

Melting point 128·5°C (with decomposition)* AHf° -199-3 kcal mol-i bb-c -188-9 kcal m o l - i A//>°(aq) Dissociation constants pKi 3-29, pK2 8-3, pK3 11-6° Reduction potential, in acid solution ca. + l-7voltsd £°(H 5 I0 6 /HI0 3 ) Crystal structure determination (HO)5IO molecules e X-ray diffraction f Neutron diffraction £ Bond lengths 1-78(2)Ä Ki=o) r(I—OH) 1-89(2) A rav(0-H) 0-96 A Infrared spectrum solid«·11 solution"»1, solid1 Raman spectrum Force constants* frd = 0) 5-4mdyneA -1 /r(I-OH) 3-OmdyneA"1 k Mass spectrum Ions:HI0 4 + ,HI0 3 + , H I 0 2 M 0 2 + , H I O + , I O + , H I + , I + a b

L. Paöesova and Z. Hauptman, Z. anorg. Chem. 325 (1963) 325. E. E. Mercer and D. T. Farrar, Canad. J. Chem. 46 (1968) 2679; J. H. Stern and J. J. Jasnosz, / . Chem. Eng. Data, 9 (1964) 534. c See Tables 81 and 82. d A. G. Sharpe, Halogen Chemistry (ed. V. Gutmann), Vol. 1, p. 13, Academic Press (1967). e Y. D. Feikema, Acta Cryst. 14 (1961) 315. f Y. D. Feikema, Acta Cryst. 20 (1966) 765. 1 H. Siebert, Z anorg. Chem. 273 (1953) 21; H. Siebert and G. Wieghardt, Z. Naturforsch. 27b (1972) 1299. h See also Table 84. 1 H. Siebert, Z. anorg. Chem. 273 (1953) 21. 1 H. Siebert, Fortschritte Chem. Forsch. 8 (1967) 470. k M. H. Studier and J. L. Huston, / . Phys. Chem. 71 (1967) 457. acid. Strongly acidic media like concentrated perchloric or sulphuric acids favour the protonated cation [I(OH) 6 ] + 6 2 5 . The principal anions present in aqueous solutions of higher pH are tetrahedrall04-,thehydrated species [(HO) 4 I0 2 ]-, [(HO) 3 I0 3 p- and [(HO)2I04P - , and the dimer [(HO^Og] 4 ~. The equilibria which link the various species (Tables 81 and 82) are apparently established rapidly and are pH-dependent. Hydration of the metaperiodate ion I04-+2H2O^H4l06is not catalysed by hydrogen ions, and is among the most rapid reactions of its kind; the specific (pseudo-first-order) rate constant is (1·9.±0·2)χ 102 sec - 1 at 20°C (ionic strength 0-1 M, pH range 3-4-5-0). The mechanism is still uncertain: both one-step and two-step paths (Scheme 14) are consistent with the kinetic data, though for various reasons the twostep mechanism (involving afive-coordinateintermediate) is considered more likely776. Solutions of periodic acid decompose slowly into iodic acid and ozonized oxygen, a reaction which is catalysed by the presence of colloidal platinum. 776 K. Kustin and E. C. Lieberman, / . Phys. Chem. 68 (1964) 3869.

pot. tit.; ultracent.; Raman u.V.; Raman

25 0-40

25 1-70

1-3 < p H < 5 · 8

pH> 7 6 < pH < 12

a

calorimetry solubility of CsI0 4 ;

20

<50

3-4 < p H

I nmr

H. Siebert and G. Wieghardt,Z. Naturforsch. 27b (1972) 1299. C. E. Crouthamel, A. M. Hayes and D . S. Martin, / . Amer.Chem. Soc. 73 (1951) 82. H. C. Mishra and M. C. R. Symons, / . Chem. Soc. (1962) 1194. S. H. Laurie, J. M. Williams and C. J. Nyman, / . Phys. Chem. 68 (1964) 1311. K. Kustin and E. C. Lieberman, / . Phys. Chem. 68 (1964) 3869. K. F. Jahr and E. Gegner, Angew. Chem., Internat. Edn. 6 (1967) 707. E. E. Mercer and D . T. Farrar, Canad. J. Chem. 46 (1968) 2679* R. M. Kren, H. W. Dodgen and C. J. Nyman, Inorg. Chem. 7 (1968) 446. J. Aveston, J. Chem. Soc. (A) (1969) 273. G. J. Buist, W. C. P. Hipperson and J. D. Lewis, / . Chem. Soc. (A) (1969) 307.

Reference

u.V., ultraviolet absorption spectroscopy; pot. tit., potentiometric titration; dil. tit., dilatometric titration; ultracent., ultracentrifuge studies.

127

kinetics dil. tit. of H 5 I0 6 with OH~

25 5-75

a q H C 1 0 4 ( < 1 0 M) 0 <; pH <; 11

Raman u.V.; pot. tit. of periodate with OHu.v. solubility of PI14ASIO4

Technique(s)a

0-70

cfl.25

Temperature (°C)

pH<0 0 <; pH <; 12

Medium

TABLE 81. INVESTIGATIONS OF PERIOD ATE EQUILIBRIA IN AQUEOUS SOLUTION

1456

CHLORINE, BROMINE, IODINE AND ASTATINE: A. J. DOWNS AND C. J. ADAMS TABLE 82. PERIODATE EQUILIBRIA IN AQUEOUS SOLUTION

Equilibrium

6-3 a 51xl0"4b 4-9x10-9 0 2-5X10-12C 40;*29 d ca. 820°

H6I06+ ^ H 5 I 0 6 + H + H5I06^H4I06-fH+ H4I06- ^ H 3 I 0 6 2 - + H + H 3 I0 6 2- ^ H 2 I 0 6 3 - + H + H4I06- ^ I 0 4 ~ + 2 H 2 0 2H 3 I0 6 2- ^ H 2 I 2 O 1 0 4 - + 2 H 2 0 a b

82.

5 d e

AH° (kcalmol - 1 )

-^298

+2-55±0-7 e ~0b + ll-9 e

H. C. Mishra and M. C. R. Symons, / . Chem. Soc. (1962) 1194. C. E. Crouthamel, A. M. Hayes and D. S. Martin,/. Amer. Chem. Soc. 73 (1951) G. J. Buist, W. C. P. Hipperson and J. D. Lewis, / . Chem. Soc. (A) (1969) 307. R. M. Kren, H. W. Dodgen and C. J. Nyman, Inorg. Chem. 7 (1968) 446. E. E. Mercer and D. T. Farrar, Canad. J. Chem. 46 (1968) 2679. (a) One-step mechanism

^μ

H OH HO

t

!<

•OH

OH

ΨΗ (b) Two-step mechanism H

o^.z -* I^*^\ 3JÖ-H

Vr·

O

L

o—it

^OH *OH

I

o HO HO

ό.Ρ-Η^ρ-Η

HO

OH

HO

OH

SCHEME 14. Mechanisms for the hydration of I 0 4 " . 624

Periodates ™ Salts of periodic acid have been isolated which contain, or purport to contain, the follow­ ing anions: metaperiodates:

I04"

mesoperiodates: orthoperiodates: triperiodates:

Ι 2 0 9 4 " , Ι 2 Οι 0 6 ", HI 2 Oi 0 5 ', H 2 I 2 Oi 0 4 ", H 3 I 2 Oi 0 3 ~ I 0 6 5 " , Η 2 Ι0 6 3 ", H 3 I0 6 2-, H4IO
together with various amounts of water of crystallization. Most cations generate several

OXYACIDS AND OXYSALTS OF THE HALOGENS

1457

compounds which are often interconvertible via aqueous solution (Scheme 13) by changing the concentration of the solution, its pH or the temperature of crystallization; some con­ versions are effected by controlled thermal dehydration, e.g. 3LiI0 4 ,3H 2 0 -> L13H4I3O14+7H2O

Definitive X-ray investigations have been few (Table 83), but extensive use has been made of infrared and Raman spectroscopy in characterizing the salts (Table 84), all of which, except the metaperiodates, involve octahedral coordination about iodine. Deriva­ tives of mesoperiodic acid ("H3IO5") invariably possess dimeric anions, as found for I2O94" (9) and Η 2 Ϊ2θ 10 4 - (10), which are formed from two I 0 6 octahedra sharing respectively a common face and a common edge.

|^(r^|^o OH (9)

O (10)

The thermal decomposition of periodates624 of various stoichiometries is accompanied by loss of water, oxygen and iodine (in some cases only one or two of these components). Reactions include partial or complete dehydration, conversion to a more stable periodate of a different composition, 5Ba(I0 4 ) 2 -> Ba 5 (I0 6 )2+4I 2 +140 2

reduction to a salt of the anions IO3 ~ or IO42 - (see below), Na 2 H 3 I0 6 -

-> N a 2 I 0 4 + 3 / 2 H 2 0 + 1 / 4 0 2 300°c

NaI04 y NaI03+*02 or formation of the metal iodide (for electropositive metals), the oxide or the metal itself. Reactions The activity of periodic acid as an oxidizing agent varies greatly as a function of pH and is capable of subtle control. In acid solution it is one of the most powerful oxidizing agents known, quantitatively and rapidly converting manganese(II) salts to permanganate, while in alkaline solution it is slightly less oxidizing than hypochlorite. Despite the ease with which they perform all the oxidations effected by iodate, periodates have not been much used in inorganic preparative and analytical operations; titration with NaI0 4 has been suggested as a method of analysing mixtures of S2 -, SO32 -, HSO3 -, S 2 0 4 2 - and S 2 0 3 2 - 777. However, there are numerous reports of the oxidation of organic compounds by periodic acid or periodates778) not merely in conjunction with perchloric acid in the "wet fire" destructive oxidation of organic compounds779j but as reagents able to execute a number of well-defined reactions. The specific cleavage of 1,2-diols (and the related compounds a-diketones, α-ketols, a-aminoalcohols and α-diamines) has been widely exploited in the realm of carbohydrates and nucleic acids, and the mechanism has been well established 777 R . L . Kaushik and R. Prosad, / . Indian Chem. Soc. 46 (1969) 405. 778 B . Sklarz, Quart. Rev. Chem. Soc. 21 (1967) 3. 779 G . F. Smith and H. Diehl, Talanta, 4 (1960) 185.

(HO) 2 I 2 0 8 4-

K4H 2 I 2 Oio,8H 2 0

l-807±0-011

l-980±0-016

1-992+0012 2017 ± 0 0 0 7

1-86

[Mg(OH 2 ) 6 ]H 3 K) 6

2-79-2-97

2-56-2-86

2-62-3-00

1-95

1-86

CdH 3 I0 6 ,3H 2 0

(HO) 3 I0 3 21-98

2-60-2-78

1-89

1-78

1-77

H5IO6

O-H · · O units

(HO) 5 IO

201

Bridging I-OI units

K42O9

I-OH units

I 2 0 9 4-

1-775 ± 0 0 0 7

Terminal I-O units

Average internuclear distances in Ä

NaI04

Compound

IO4-

Unit

TABLE 83. AVERAGE INTERNUCLEAR DISTANCES IN PERIODATES

A. Kaiman and D . W. J. Cruickshank, Acta Cryst. B26 (1970) 1782. B. Brehler, H. Jacobi and H. Siebert, Z . anorg. Chem. 362 (1968) 301. Y. D . Feikema, Acta Cryst. 20 (1966) 765. A. Braibanti, A. Tiripicchio, F . Bigoli and M. A. Pellinghelli, Acta Cryst. B26 (1970) 1069. F. Bigoli, A. M. M. Lanfredi, A. Tiripicchio and M. T. Camellini, Acta Cryst. B26 (1970) 1075. H. Siebert and H. Wedemeyer, Angew. Chem., Internat. Edn. 4 (1965) 523; A. Ferrari, A. Braibanti and A. Tiri­ picchio, Acta Cryst. 19 (1965) 629.

References

OXYACIDS AND OXYSALTS OF THE HALOGENS

1459

TABLE 84. CHARACTERISTIC VIBRATIONAL FREQUENCIES OF PERIODATES*

Frequency (cm - 1 )

Character of vibration

3100-3400 2200-2900 1600, 3400 1070-1310 840-950 600-850 500-600 250-500

O-H stretching: weakly hydrogen-bonded O-H stretching: strongly hydrogen-bonded Water of crystallization I-OH deformation I-OH torsion (found only with strong hydrogen bonds) I-O stretching I-O-I stretching OIO deformation

a H. Siebert, Fortschritte Chem. Forsch. 8 (1967) 470; H. Siebert and G. Wieghardt, Spectrochim. Acta, 27A (1971) 1677; Z. Naturforsch. 27b (1972) 1299; see also Table 76.

(Scheme 15). In rigid systems only cw-functional groups are oxidized, the specificity being due to the cyclic intermediate o ""O^ll HO-CMe, O ^ l l HO-CMe,2 O

O

0

"O. II .OCMe,2 ^Ι^ | H(T 1 ^O-CMc,2 OH

^-.O-CMe, _

v

xrii o:.

i

O

Me

—

o

IIII JOCMet, JOCMcu

Ο-ΙΓ

I ^

II "^O-C-OH KM II O

Me

-♦

Me2CO Me2CO

Me,CO Me „CO

J>

MeC^ OH

+

10;

+

IO;

H20

3

SCHEME 15. Periodate fission of glycols and a-ketols.

Periodate has also been applied to the oxidation of phenols (to quinones), sulphides (to sulphoxides) and hydrazine derivatives: R3CNHNH2 X-^U R3CH+N2

Periodate Complexes and Heteropolyanions By mixing solutions of alkali-metal periodates with solutions of appropriate transition metal salts (accompanied, if necessary, by chemical or electrochemical oxidation), it has been possible to isolate salts containing the periodate complexes and heteropolyanions listed in Table 85; contact with a protonic cation-exchange resin affords, in some cases, the free acid in solution, e.g. H2[FeI06] and Η 5 [ϊ(Μο0 4 ) 6 ]. The iodine and metal atoms in these compounds are linked via I-O-M bridges; the iodine atom is invariably surrounded by an octahedron of oxygens, but the coordination of the metal atom is determined by the crystalfield and other demands peculiar to its oxidation state. The periodate complexes exhibit high formation constants, e.g. [Cu(I06)2]7 "* ca. 1010 and [Co(I06)2]7 ~ ca. 1018, and may contain metal atoms in unusually high formal oxidation states, e.g. Cu(III) and Ag(III); they may be broken down in strongly acid solution. Like the

1460

CHLORINE, BROMINE, IODINE AND ΑβΤΑΉΝΕΙ A. J. DOWNS AND C. J. ADAMS TABLE 85. CONSTITUTION OF PERIODATE COMPLEXES AND HETEROPOLYANIONS»

Metal

n

Anion Periodate complexes [M4006)3]ft[MI06]n-

3 1 2 6 7 11

[M(I06)2]*-

[Μσο6)3]η-

Heteropolyanions [I(M04)6]n- n [I02(M04)4] [I0 5 (M0 4 )]-

5 5 5

Fe(III), Co(III) Ni(IV), Mn(IV) Fe(IH), Co(III) Pd(IV), Pt(IV), Ce(IV) Co(III), Fe(III), Cu(III), Ag(III), Au(III) Mn(IV) Mo(VI), W(VI) Mo(VI) Mo(VI), W(VI)

»For references see: H. JSiebert, Fortschritte Chem. Forsch. 8 (1967) 470; M. Drätovsky and L. Paöesovä, Russ. Chem. Rev. 37 (1968) 243.

simple periodates, the complex anions often contain protons (attached to I-O bonds) which ionize only with great difficulty if at all: the green acid H7[Co(I06)2] shows only five protons which can be titrated (ρΚχ, small; pK2 = 1-95; pK3 = 7-1; pK4 = 8-0; pK5 = 12-1)780, while Hn[Mn(I06)3] is apparently heptabasic (ρΚ^ and pK2 ^ 0; pK3 = 2-75; pK4 = 4-35; pK5 = 5-45; pK6 = 9-55; pK7 = 10-45)780. Crystal-structure determinations show the salts Na3KH3[Cu(I06)2],14H2078i and Na7H4[Mn(I06)3],17H20782 to contain the anions (11) and (12), in both of which the periodate functions as a bidentate ligand. The metal enjoys square-planar coordination in the diamagnetic copper salt and octahedral geometry in the paramagnetic manganese(IV) compound. The heteropolyanions (Table 85) have received comparatively little attention. o o H,

I

.ο^

^χκ

^ic

' <

Ψ*

o

ο"Ίχο

o

o x l ^,ο o

(")

(12)

(F) P A R A M A G N E T I C O X Y A N I O N S

The action of ionizing radiation on chlorates and bromates produces defect centres, some of which have been identified spectroscopically at low temperatures as radical anions; these include [Cl-C10 2 ]- and C1032"in y-irradiated KC1036io, [Br-Br0 3 ]-, [BrO-Br0 3 ]780 M . W. Lister, Canad. J. Chem. 39 (1961) 2330. 78i I. Hadinec, L. JenSovsky, A. Linek and V. Synecek, Naturwiss. 47 (1960) 377. 782 A . Linek, Czech. J. Phys. 13 (1963) 398. 783 R. c . Catton and M. C. R. Symons, Chem. Comm. (1968) 1472.

OXYACIDS AND OXYSALTS OF THE HALOGENS 2

1461

726

and Br0 3 in KBr0 3 treated with thermal neutrons , and CIOH" in y-irradiated BaCl2,2H20783. While these transient species trapped in crystalline lattices are interesting intermediates in the radiolytic decomposition of the halogen oxyanions, macroscopic quantities of several paramagnetic materials have been isolated which are apparently derivatives of iodine(VI)624. White or yellow solids with the stoichiometries M I 2 I0 4 (M1 = Li or Na), Be3l209,4H20, M n I0 4 ,nH 2 0 (n = 1, M = Ca or Ba; n = 0-5, M = Mg; n = 0, M = Pb) and A12(I207)3,2H20 are formed on heating certain hydrated periodates to ca. 200°C, and themselves decompose at ca. 400°C. Ba 2 l20 9 ,3H 2 0

220°O

> 2BaI04,H20+H20+i02

The amorphous solids are stable in air at room temperature, but hydrolyse in water to give iodine(V) and iodine(VII) anions. Their mass magnetic susceptibilities lie in the range 0-2-5-2 x 10~6 cgs units, giving moments in the range 0-5-2-1 B.M. There is no chemical evidence for the presence of peroxides, and the possibility that impurities are responsible for the paramagnetism has apparently been excluded624. It is claimed that crystalline barium iodate(VI) can be obtained by heating the amorphous BaI04,H 2 0 in a eutectic mixture of sodium and potassium nitrates at 250°C; the solid gives rise not only to a sharp X-ray diffraction pattern quite different from that of barium iodate(V), but also to a measurable esr signal784. Tentatively identified by their ultraviolet-visible absorption spectra, the ion IO42 - and the radical -I0 4 are believed to be formed as transients during the pulse radiolysis of aqueous periodate785. Up to the present time there has been no report of long-lived oxy-chlorine or oxybromine salts analogous to the iodate(VI) derivatives, though the parent oxide CIO3 can be prepared and handled on the macroscopic scale (see p. 1372). However, there is an ostensible likeness between iodate(VI) and the oxy-selenium(V) and oxy-tellurium(V) compounds which appear to be formed in the thermal decomposition of Se0 3 , Te0 3 , Te(OH)o and certain tellurates624. (G) A N A L Y T I C A L C H E M I S T R Y OF THE OXYANIONS572,786-789

The limitations of space preclude any attempt to itemize the many methods which have been used or suggested to detect, separate and estimate the oxyanions of the halogens. The intention is rather to describe the way in which the chemical properties of these species have been turned to advantage in their quantitative analysis. No references will be given for procedures which are already well-documented in standard texts on analytical chemistry786»787, or in previous reviews of the analysis of oxyanions formed by chlorine788, bromine788»789 and iodine78?; the reader is also referred to these sources for details of empirical, qualitative tests for the presence of the anions. Discussion is restricted to methods in current use. 784 M . Drätovsky, J. Juläk and V. Trkal, Coll. Czech. Chem. Comm. 32 (1967) 3977. 785 F . Barat, L. Gilles, B. Hickel and B. Lesigne, Chem. Comm. (1971) 847. 786 A. I. Vogel, A Text-book of Quantitative Inorganic Analysis, 3rd edn., Longmans, London (1961). 787 L . A. Haddock, Comprehensive Analytical Chemistry (ed. C. L. Wilson and D. W. Wilson), Vol. IC, pp. 341-366, Elsevier (1962). 788 L . A. Haddock, Mellows Comprehensive Treatise on Inorganic and Theoretical Chemistry, Supplement II, Part I, pp. 660-685, 800-812, 937-956, Longmans, London (1956). 789 p. G. W. Scott and M. L. Parker, Bromine and its Compounds (ed. Z. E. Jolles), pp. 768-772, Benn, London (1966).

1462

CHLORINE, BROMINE, IODINE AND ASTATINE: A. J. DOWNS AND C. J. ADAMS

The methods outlined here may also be applied to other "positive halogen" derivatives (e.g. interhalogens, halogen esters of oxyacids, etc.), since on hydrolysis in aqueous solution these compounds commonly yield oxyanions. Apart from gravimetric and spectrophotometric techniques, the methods used in determining the oxyhalogen anions involve their reduction. Polarographic analysis has been achieved for all the anions except C104 - and Br0 4 ~ (Table 86), and, in principle, a wide variety of chemical reducing agents could also be used; the literature abounds with sugges­ tions. In practice, however, only a small number of reductants is routinely employed, and only three methods are of major significance: (1) Direct titration with a reducing agent, the end-point being located with indicators, or by potentiometry, amperometry, coulometry or conductimetric measurements. It is plainly important for titrimetry that the reaction be rapid (although back-titration of excess reductant is always possible), so that, for less reactive species like C103 ~, acid solution and a catalyst may be necessary. Sodium arsenite reduces all the anions except C104 ~ and Br0 4 ~, and is by far the best titrant (see below); starch-potassium iodide is a common chemical indicator, and back-titration, when necessary, is conveniently carried out with standard bromate solution. (2) Reduction to halide, which is subsequently determined by a standard procedure (p. 1331), is a method open to all the oxyanions, but is used especially for the chlorine species. The strength of the reducing agent and the reaction conditions needed depend on the anion being reduced. Perchlorate requires vigorous methods such as fusion or reduction with Ti 3 + ; other reagents which have been used include stannous chloride, ferrous sulphate and sulphurous acid. (3) With the exception of perchlorate, all the anions (and CIO2) have been estimated by analysis of the iodine which they liberate (directly or indirectly) from acidified potassium iodide solution. Reaction of hypochlorite with iodide is rapid at pH 8, but chlorates and bromates in strongly.acidic solution require catalysis (by ferrous salts and molybdates respectively) for a speedy reaction. In some cases the Dietz method may be used, whereby reaction of the oxyanion with bromide in strongly acid solution liberates bromine, which is subsequently allowed to react with iodide ions. The iodine has been determined spectrophotometrically as I 3 ~ or by titration, for example, with thiosulphate. Other reduction procedures are discussed below in connection with the analysis of mixtures of the oxyanions; gravimetric and physical methods of analysis are listed in Table 86. Because C104 ~ is so hard to reduce at room temperature, the analysis of perchlorates has presented some difficulties. Many workers have reduced perchlorate to chloride under forcing conditions (e.g. by fusion with sodium carbonate in a platinum crucible or by com­ bustion with an ammonium salt) and subsequently estimated the chloride formed. Reducwith an iron(III) solution), but the method requires an inert atmosphere; reaction with vanadium(II) sulphate produces vanadium(IV), which may be determined spectrophotometrically. Recent developments have involved gravimetric analysis by precipitation of insoluble perchlorates using large organic cations, e.g. Pr^M4" (M = P, As or Sb), though other large anions are liable to interfere; potentiometric, conductimetric and amperometric titration of perchlorate solutions using these reagents has also been proposed. Colonmetric reagents (mainly complicated quaternary ammonium salts such as methylene blue)

1463

OXYACIDS AND OXYSALTS OF THE HALOGENS TABLE 86. SPECTROPHOTÖMETRIC, GRAVIMETRIC AND POLAROGRAPHIC METHODS OF ANALYSING OXYHALOGEN SPECIES

Spectrophotometry Oxyhalogen species

ciocio 2 C10 2 C103C104BrO~ Br02" Br03Br04" IOIO3IO4-

Direct At At At At

290 m/*i 261 m/*i 250 mji2 360 m/ja

With a chromophoric agent

Gravimetric analysis

23 24,25 4 5 6-10

19-21

11-13 14 At 222-5 m/x3

Polarography

15 16-18

22

24,25 26 27,28 27-30 27,31 32 31 33

1 T. Chen, Anal, Chem, 39 (1967) 804. HCIO2 and C10 2 ~ have the same extinction coefficient at 250 τημ; C. C. Hong and W. H. Rapson, Canad.J. Chem, 46 (1968) 2061. 3 G. O. Aspinall and R. J. Ferrier, Chem, andlnd, (1957) 1216. 4 With tyrosine; P. Kerenyi and P. Kuba, Chem, Zveste, 17 (1963) 146. 5 With o-tolidine; P. Urone and E. Bonde, Anal, Chem, 32 (1960) 1666. <> Methylene blue-perchlorate complex extracted into C1CH 2 CH 2 C1; I. Swasaki, S. Utsumi and C. Kang, Bull, Chem, Soc, Japan, 36 (1963) 325; complex precipitated; G. M. Nobar and C. R. Ramachandran, Anal, Chem, 31 (1959) 263. 7 Crystal violet-perchlorate complex extracted into C 6 H 5 C1; S. Uchikawa, Bull, Chem, Soc, Japan, 40 (1966) 798. 8 Ferroin perchlorate extracted into C3H7CN; J. Fritz, J. E. Abbink and P. A. Campbell, Anal, Chem, 36 (1964) 2123; ferroin perchlorate precipitated; Z. Gregorowicz, F . Bull and Z. Klima, Mikrochim, Ichnoanal, Acta (1963) 116. 9 With α-furildioxime; N . L. Trautwein and J. C. Guyon, Anal, Chem, 40 (1968) 639. 10 Neutral red-perchlorate complex extracted into C 6 H 5 N 0 2 ; M. Tsubouchi and Y. Yuroko, Bunseki Kagaku, 19 (1970) 966. 11 With σ-arsanilic acid; J. C. MacDonald and J. H. Yoe, Anal, Chim, Acta, 28 (1963) 383. 12 With o-aminobenzoic acid; M. H. Hashmi, H. Ahmad, A. Rashid and A. A. Ayaz, Anal, Chem, 36 (1964) 2028. " W i t h l,2,3-tris-(2-diethylaminoethoxy)benzene; I. Odler, Anal, Chem, 41 (1969)116. 14 Crystal violet-perbromate complex extracted into C6H5CI; L. C. Brown and G. E. Boyd, Anal, Chem, 42 (1970) 291. is With 1-naphthylamine; A. S. Vorob'ev, Uch, Zap, Udmurtsk, Gos, Bed, Inst, 11 (1957) 150; Chem. Abs. 54 (1960) 6386d. 16 Crystal violet-periodate complex extracted into benzene; C. E. Hedrick and D . A. Berger, Anal. Chem, 38 (1966) 791. " With benzhydrazide; A. M. Escarilla, P. F . Maloney and P. M. Maloney, Anal, Chim, Acta, 45 (1969) 199. i* With o-dianisidine; M. Guernet, Compt, rend. 254 (1962) 3688. is» As [(n-C5Hu)4N]C10 4 ; R. G. Dosch, Anal. Chem. 40 (1968) 829; [(n-C 5 Hn)4N]Br used in the potentiometric titration of CIO4". 20 As [Ph4P]C104 or [Ph 4 Sb]C10 4 ; H . H. Willard and L. R. Perkins, Anal. Chem. 25 (1953) 1634; [Ph4Sb] 2 S0 4 used in the amperometric titration of CIO4-; M. D . Morris, Anal, Chem, 37 (1965) 977. 2

1464

CHLORINE, BROMINE, IODINE AND ASTATINE: A. J. DOWNS AND C. J. ADAMS

Table 86 (cont.) 21 As [Ph4As]C104; T. Okuba, A. Fumio and T. Teraoka, Nippon Kagaku Zasshi, 89 (1968) 432; [PhtAsJCl used in amperometric and conductimetric titrations of CIO4-; G. M. Smith, Ind. Eng. Chem.y Anal. Edn. 11 (1939) 186; R. J.Baczukand W. T. Bolleter, Anal. Chem. 39 (1967) 93. 22 As [Ph4As]I04; S. H. Laurie, J. M. Williams and C. J. Nyman, / . Phys. Chem. 68(1964)1311. 23 E. N. Jenkins, / . Chem. Soc. (1951) 2627. 24 O. Schwarzer and R. Landsberg, / . Electroanal. Chem. 14 (1967) 339. 25 G. Raspi and F. Pergola, / . Electroanal. Chem. 20 (1969) 419. 26 Requires OSO4 as a catalyst; L. Meites and H. Fofsass, Anal. Chem. 31 (1959) 119. 27 H. Fuchs and R. Landsberg, Anal. Chim. Acta, 45 (1969) 505. 28 F. Pergola, G. Raspi and A. Massagei, / . Electroanal. Chem. 22 (1969) 175. 29 A. F. Krivis and G. R. Supp, Anal. Chem. 40 (1968) 2063. 30 N. Velghe and A. Claeys, Bull. Soc. Chim. Beiges, 75 (1966) 498. 3i J. Proszt, V. Cieleszby and K. Gyorbiro, Polarographie, p. 332, Akademie Kiado, Budapest (1967). 32 O. Souchay, Anal. Chim. Acta, 2 (1948) 17. 33 R. D. Corlett, W. C. Breck and G. W. Hay, Canad. J. Chem. 48 (1970) 2474.

have also been used, but most require a precipitation or extraction step before measurement, and the method is subject to interference from other ions. In the absence of interfering species, the determination of individual anions is straight­ forward using any of the methods described in the foregoing paragraphs or listed in Table 86. However, the successful analysis of mixtures of the anions may require considerable ingenuity. In some cases, differences in reaction rates may be exploited, and several species may be successively titrated by careful pH control and the use of catalysts. It may be possible selectively to reduce or otherwise destroy some components of the mixture, so that a series of determinations of the total oxidizing power may be made with different com­ ponents rendered inactive. Again, a combination of several techniques may be needed. The remainder of this section is devoted to some specific problems of this sort and to some of the solutions which have been proposed. ClO -, C102 -, C/02, C103 - and C104 ~ Hypochlorite, chlorite and chlorate may be determined by potentiometric titration with arsenite in, respectively, alkaline solution, alkaline solution with OSO4 as a catalyst, and acid solution with Os0 4 as a catalyst; a similar procedure but using amperometric titrations has been described for mixtures containing Cl2, chloramine, C102 and C102~. Recent investigations of these systems have utilized spectroscopic procedures. Thus, C102, having been determined spectrophotometrically at 360 m/*, may be removed by passing an inert gas through the solution (simultaneously sweeping out some HOC1), whereupon chlorite and chlorous acid can be estimated together at 250 m/x, where they have the same absorbance; the total oxyanion concentration is determined iodometrically790. Alternative procedures involve the spectrophotometric measurement of I3 - for CIO - and (C10 2 - + C10-) (following oxidation of KI at pH 8-3 and pH 2-3 respectively) and of Fe3 + for chlorate (following oxidation of ferrous sulphate)791»792. In either case, perchlorate may be determined as the difference between the chloride formed on total reduction of the mixture, e.g. with Ti3 +, and that formed on partial reduction of the mixture, e.g. with hot 790 c . C. Hong and W. H. Rapson, Canad. J.Chem. 46 (1968) 2061. 791 T. Chen, Anal. Chem. 39 (1967) 804. 792 L. A. Prince, Anal. Chem. 36 (1964) 613.

OXYACIDS AND OXYSALTS OF THE HALOGENS

1465

H2SO3, which reduces only chlorates, chlorites and hypochlorites. The presence of C102 interferes with iodometric measurements since it too oxidizes I~ to I2; in analysing such solutions, the iodine liberated at pH 8-3 and pH 2-3 corresponds to (QO~+ I/2CIO2) and (ClO -+2C10 2 ~+5/2C102) respectively. Mixtures Containing BrO~9 Br02~, BrO$~ and Br04~~ The +1, +3 and +5 oxidation states of bromine are readily determined by potentiometric titration with sodium arsenite: hypobromite is so determined in cold alkaline solution, bromite in alkaline solution with Os0 4 as a catalyst, and bromate in acid solu­ tion^3»794. Hypobromite may be determined in the presence of bromite by potentiometric titration with ferrocyanide794. The colorimetric estimation of Br0 4 ~ with crystal violet is apparently insensitive to the presence of other oxy-bromine anions. Mixtures Containing 10~, 101~ andPeriodates Hypoiodites may be measured in the presence of other oxy-iodine species by their reaction with phenol which produces a precipitate of 2,4,6-triiodophenol. Two portions of the solution are acidified and one is treated with phenol; both are then treated with potas­ sium iodide, and the difference in the amounts of iodine liberated corresponds to the hypoiodite content of the solution. Selective iodometry allows iodates and periodates to be estimated: the iodine liberated from KI at pH 4-4-7 corresponds to iodine(VII)'s being reduced to iodine(V), but at lower pH iodate is itself reduced with the liberation of further iodine. Similarly, in a solution buffered with sodium carbonate, periodate is reduced only to iodate. Periodate is stabilized against reaction with I - in the presence of molybdate, when a heteropolyanion is formed; under these conditions iodate liberates iodine, which may be estimated with thiosulphate, and periodate is then regenerated by treatment with oxalic acid795.

Chlorates, Bromates and Iodates These species can be separated by ion-exchange chromatography on a strongly basic resin796, by paper electrophoresis, and by paper chromatography; their characteristic infrared spectra have been used as a basis for quantitative determinations797. Chemical methods for their estimation in the presence of one another depend on the increase in reactivity as the atomic number of the halogen increases. Thus, in dilute acid in the presence of chlorate, bromate is reduced to Br~ by oxalic acid, and oxidizes I - to I2; chlorate undergoes neither reaction. In a very weakly acid solution containing Br0 3 -, ΙΟ3 - and F ~, only IO3- reacts with I~. Alternatively, a simple colorimetric procedure798 for analysing Br0 3 ~ and IO3 - individually or in admixture uses a 1:1 mixture of isonicotinic acid hydrazide and 2,3,5-triphenyltetrazolium chloride, which in dilute hydrochloric acid gives a pink colour with IO3- at 20°C, whereas the bromate colour develops only on heating; chlorate apparently does not interfere. Polarography can also distinguish Br0 3 - or IO3 ~ in the presence of CIO3 ~, since in the absence of a catalyst the oxy-chlorine anion is only slowly reduced. 793 p. Norkus and S. Stulgiene, Zhur. Anal. Khim. 24 (1969) 1565. 794 T . Andersen and H. E. L. Madsen, Anal. Chem. 37 (1965) 49. 795 R. Belcher and A. Townshend, Anal. Chim. Acta, 41 (1968) 395. 796 M . Kikindai-Cassel, Ann. Chim. 3 (13) (1958) 5. 797 M . W. Miller, R. H. Philp, jun. and A. L. Underwood, Talanta, 10 (1963) 763. 798 M . H. Hashmi, H. Ahmad, A. Rashid and F. Azam, Anal. Chem. 36 (1964) 2471.

1466

CHLORINE, BROMINE, IODINE AND ASTATINE: A. J. DOWNS AND C. J. ADAMS 5. H A L O G E N D E R I V A T I V E S OF O X Y A C I D S

Introduction

It has been known for some forty years that in polar non-aqueous solvents chlorine, bromine and iodine undergo reactions with silver salts 7 ", e.g. Cl2+AgN03

EtOH

► AgCl+CINO3

That the resulting solutions conducted electricity and liberated the elemental halogen at the cathode during electrolysis convinced earlier researchers that the reactions could be des­ cribed as metatheses, and the products as salts of halogen cations, e.g. Cl+Cl-+Ag+N03- ->Ag+Cl+Cl+N03-

It proved impossible to isolate the pure halogen derivatives from these solutions, although adducts with organic bases, e.g. C5H5N,IOCOCH3 and [C1(C8H11N)2]N03, were readily obtained; the chemistry of these adducts was discussed in Section 4A (p. 1345). More recent synthetic work has shown the pure halogen "salts" to be covalent compounds with halogenoxygen bonds, and it is the chemistry of these materials—formally esters—with which the present section is concerned. Halogen(I) nitrates, fluorosulphates, perchlorates and carboxylates have structures related to that of the parent acid, the hydroxylic hydrogen being replaced by fluorine, chlorine, bromine or iodine. The thermal stability of these compounds decreases as the atomic number of the halogen increases, although the estimation of stabilities is complicated by the susceptibility of the materials to decomposition in the presence of trace impurities. In common with esters of the hypohalous acids (which have similar R-O-Hal skeletons), they add across olefinic double bonds with rupture of the oxygen-halogen linkage: Bu^OCl

Bu'OCHrCHzCl «

CH2=CH2

ClONO-

> C1CH2CH20N02

They are distinguished from covalent hypochlorites by their reactions with metal halides: «

F5SOCI

H

CsF

CIOS0 2F

► CsS0 3 F+ClF

That comparable hypochlorites, like F5SOCl and CF3OCl, fail to react with CsF is attrib­ uted to the instability of the anion formed on heterolysis of the Cl-O bond. Further indication of the "positive" character of the halogen in oxyacid derivatives is the replacement of chlorine in reactions with bromine or iodine, e.g. 2CIOCIO3+Br 2 -► 2BrOC10 3 +Cl 2

No chlorine(III) oxyacid derivatives are known, but bromine and iodine are readily converted to the tripositive state, even by silver salts in non-aqueous solution: 2Br 2 +3 A g N 0 3 -> Br(N0 3 ) 3 +3 AgBr

The trivalent bromine and iodine compounds are usually more stable than the correspond­ ing monovalent species (whose instability is comparable with the ready disproportionation of BrF and IF). Few derivatives of this kind containing pentavalent halogen atoms have been reported, and none is known for the heptavalent elements. 799 M . I. Ushakov, / . Gen. Chem. (U.S.S.R.) 1 (1931) 1258.

1467

HALOGEN DERIVATIVES OF OXYACIDS

Halogen Fluorosulphates Fluorosulphate derivatives of the halogens have featured in a recent general review of fluorosulphates800. The most important compounds are: BrOS0 2 F Br(OS0 2 F) 3 M I [Br(OS0 2 F) 4 ]

C10S0 2 F

Br 3 OS0 2 F(?) C10 2 OS0 2 F

IOS0 2 F I(OS0 2 F) 3 M^COSOzFW ICI 2 OS0 2 F I 3 OS0 2 F I0 2 OS0 2 F IF 3 (OS0 2 F) 2

That much of the chemistry of interhalogen compounds—exchange, addition, displacement and complexation reactions—can be simulated by these materials is evident from the reactions of the iodine fluorosulphates. Thus, in many respects, the fluorosulphate group is behaving as a pseudohalogen (see p. 1122). FOS0 2 F

K1C1,

I F 3( O S O 2F ) 2

S206F2

I,OSO,F

C12

sr

^gz

I(OSOJF),

SO,

K[!(OS0 2 F) 4 ]

I205(or KIO3) SCHEME 16. Iodine fluorosulphates.

Most of the halogen fluorosulphates are highly sensitive to moisture, and decompose slowly on standing at room temperature. The simpler compounds are formed by mixing the halogen and peroxydisulphuryl difluoride, FSO2OOSO2F, in requisite proportions under appropriate conditions. The equilibria established when various proportions of halogen and peroxydisulphuryl difluoride are dissolved in fluorosulphuric acid (or related solvents) have been studied spectroscopically, conductimetrically and by measurements of colligative properties; this topic has been alluded to in Section 4A (p. 1341) in connection with the formation and characterization of halogen cations. Halogen(I) Fluorosulphates Direct combination of the halogen with an equimolar quantity of peroxydisulphuryl difluoride affords C10S0 2 F, BrOS0 2 F or IOS0 2 F (Table 87). C10S0 2 F may also be pre­ pared in high yield from sulphur trioxide and chlorine monofluoride801, while BrOS0 2 F and 800 A. A. Woolf, New Pathways in Inorganic Chemistry (ed. E. A. V.Ebsworth, A.G.Maddock and A. G. Sharpe), p. 327, Cambridge (1968). sol W. P. Gilbreath and G. H. Cady, Inorg. Chem. 2 (1963) 496.

-36-3

liquid* liquid11

gasf solidf liquid8 liquid11 2-6f -33-9

gas®

-41-3

liquid*

black* 51-5d

yellowb -84-3b 45.p 7-66b 24-0*

colourless6 -158-5° -31·3β 5·35β 22·1 β

red-brownc — 31-5° 117-3° 9.94c 25-8c

20-60°C d

0-50°C c

125°C, pressure1*

a

to

Prepared from SO3 and excess F2. W. P. Gilbreath and G. H. Cady, Inorg. Chem. 2 (1963) 496. c F. Aubke and R. J. Gillespie, Inorg. Chem. 7 (1968) 599. d F. Aubke and G. H. Cady, Inorg. Chem. 4 (1965) 269. β F. B. Dudley, G. H. Cady and D. F. Eggers, jun., / . Amer. Chem. Soc. 78 (1956) 290. r K. O. Christe, C. J. Schack and E. C. Curtis, Spectrochim. Acta, 26A (1970) 2367. g A. M. Qureshi, L. E. Levchuk and F. Aubke, Canad. J. Chem. 49 (1971) 2544; F. Aubke and A. M. Qureshi, Inorg. Chem. 10 (1971) 1116. h F. A. Hohorst and J. M. Shreeve, Inorg. Chem. 5 (1966) 2069.

a

Raman spectrum !9F nmr spectrum Force constant,/ r (0-X) (mdyne A" 1 ) 19F chemical shift, S(S-F) (ppm relative to CCl 3 F) h

Reaction conditions, X2+S2O6F2 -> 2XOS0 2 F Physical properties Colour Melting point (°C) Boiling point (°C) A# v a p (kcalmol-i) Trouton's constant (cal deg"1 mol" 1 ) Molecular spectra and parameters Infrared spectrum

IOSO2F

BrOS0 2 F

CIOSO2F

FOS0 2 F

TABLE 87. PROPERTIES OF HALOGEN(I) FLUOROSULPHATES

HALOGEN DERIVATIVES OF OXYACIDS

1469

IOSO2F are formed in the thermal decomposition of the respective trisfluorosulphates. The vibrational spectrum of CIOSO2F is consistent with a molecular structure of Cs symmetry (analogous to those of HOSO2F and FOSO2F) in the solid, liquid and vapour phases. BrOSC>2F is a viscous liquid which somewhat resembles BrF3; the specific conductivity (7-21 x 10~4 ohm - 1 cm - 1 at 25°C) signifies some self-ionization, possibly according to the equation802 3BrOS0 2 F ^ [Br 2 OS0 2 F] + + [Br(OS0 2 F) 2 ]-

Influorosulphuricacid solution IOS0 2 F disproportionates803: 5IOS0 2 F ^ 2 I 2 + + I ( O S 0 2 F ) 3 + 2 S 0 3 F -

Disproportionation of BrOSC^F occurs only in more acidic media, e.g. the "super-acid" HS0 3 F-SbF 5 -S0 3 S04. Halogen(I) fluorosulphates are extremely reactive compounds. Organic materials (including Kel-F) are attacked at ambient temperatures, and hydrolysis of CIOSO2F proceeds violently with evolution of some oxygen. Photolysis of CIOSO2F at low tempera­ ture produces Cl2 and FS0 2 OOS0 2 F, while the ready addition of XOS0 2 F (X = Cl, Br or I) across olefinic double bonds also involves rupture of the halogen-oxygen bond. CF 2 =CF 2 +BrOS0 2 F - * CF 2 BrCF 2 OS0 2 F

BrOSC^F reacts with molecular chlorides according to the general equation MC^+^BrOSOzF -+ MC1, -y(OS02F)y+>>BrCl

Compounds thus derived include [C(0)OS02F]2, C(OS02F)4 and Cl2P(0)OS02F8°5. C10S0 2 F similarly attacks alkali-metal salts to produce chlorine(I) derivatives of the anion together with the alkali-metal fluorosulphate: M IC 1 0 4

ClOClOs+MPSOaF <

C10S0 2 F

MXF

> ClF+M^OaF

Halogen(III) Fluorosulphates Whereas IF3 readily disproportionates to iodine and IF5, both the trisfluorosulphates Br(OS02F)3 and I(OS02F)3 (prepared from the halogen and an excess of FS0 2 OOS0 2 F) are in this respect quite stable at room temperature. At 80°C I(OSC>2F)3 decomposes slowly in vacuo to give IOS02F, S0 3 , IF(OS02F)2 and possibly I(OS02F)5 80<>. The Raman spectra of the pale yellow solids Br(OS02F)3 (m.p. 59°C) and I(OS02F)3 (m.p. 32°C) contain lines attributable to both terminal and bridgingfluorosulphategroups807. Ampholytic behaviour is noted influorosulphuricacid solution, wherein cryoscopic and conductimetric measure­ ments have identified the equilibria808 I(OS0 2 F) 3 ^ [I(OS0 2 F) 2 ] + +SO3F-; I ( O S 0 2 F ) 3 + S 0 3 F - ^ [I(OS0 2 F) 4 ]-;

Kb ~ 10"5 mol kg"i Ka ~ 10 mol~i kg

802 F . Aubke and R. J. Gillespie, Inorg. Chem. 7 (1968) 599. 803 R . j . Gillespie and J. B. Milne, Inorg. Chem. 5 (1966) 1577. 804 R . j . Gillespie and M. J. Morton, Chem. Comm. (1968) 1565. 805 D. D. Des Marteau, Inorg. Chem. 7 (1968) 434. 806 F. Aubke and G. H. Cady, Inorg. Chem. 4 (1965) 269. 807 H. A. Carter, S. P. L. Jones and F. Aubke, Inorg. Chem. 9 (1970) 2485. 808 R. j . Gillespie and J. B. Milne, Inorg. Chem. 5 (1966) 1236.

1470

CHLORINE, BROMINE, IODINE AND ASTATINE: A. J. DOWNS AND C. J. ADAMS

The Raman spectra of the ions [Br(OS02F)4] - and [I(OS02F)4] - in alkali-metal salts imply square-planar M 0 4 skeletons (compare BrF 4 ~ and IC14_), as well as suggesting that the Br-0 bonds involve a greater degree of covalency than the 1-0 bonds807. Oxidation of IOS0 2 F with chlorine gives a product of approximate composition IC1 2 (0S0 2 F), while reaction between stoichiometric amounts of I2 and IOS0 2 F affords the dark brown, hygroscopic solid I3OSO2F, which melts at 92°C with the liberation of iodine and forms the cation I 3 + on dissolution in fluorosulphuric acid806. B^OSC^F has not been isolated as such, but its probable presence as an impurity in "BrOSC^F" prepared from FS0 2 OOS0 2 F and a slight excess of bromine is indicated802»809 by discrepancies between the appearance, physical properties and reactivity of such samples and those reported more recently for pure specimens. In the "super-acid" solvent HS0 3 F-SbF 5 -3S03, 3:1 mixtures of Br2 and FS0 2 OOS0 2 F produce Br 3 + and SO3F"*810. Halogen(V) Fluorosulphates While there is no evidence for the direct formation of I(OS0 2 F) 5 from iodine and a large excess of FSO2OOSO2F, the compound is possibly one of the products of the thermal decomposition of I(OS0 2 F) 3 . Two oxyhalogen(V) fluorosulphates have been reported, viz. CIO2SO3F810 and IO2SO3F646'805; these have already been cited in the general context of cationic oxyhalogen species (p. 1352). Halogen Nitrates The established nitrate derivatives of chlorine, bromine and iodine are CIONO2

BrON02 MI[Br(ON02)2] Br(ON02)3 Br02N03

IONO2 MI[I(ON02)2] I(ON02)3 MTOONO^]

The chemistry of these compounds as known up to 1961 has been reviewed by Schmeisser and Brändle811, while the status of bromine nitrates (up to 1966)812 and of chlorine nitrate (up to 1968)572 has been the burden of more recent reviews. The action of silver nitrate on an alcoholic solution of chlorine, bromine or iodine gives rise to the halogen(I) nitrate; with excess silver nitrate, bromine and iodine give Br(N03)3 and Ι(Ν0 3 ) 3 respectively. While the pure materials cannot be isolated from these solutions, adducts with organic nitrogen bases, e.g. [(C5H5N)2Br]N03, are readily formed. Electrolysis of alcoholic solutions of INO3 releases iodine at the cathode. Chlorine(I) nitrate is a product of almost any reaction between a chlorine oxide and an oxide of nitrogen, being formed by the combination of ClO and NO2 radicals. Preparatively, the best synthesis is depicted by the equation580 C120+N205 - ^ 2C10N02 809 j . E . Roberts a n d G . H . Cady, / . Amer. Chem. Soc. 82 (1960) 352. 810 R . j . Gillespie a n d M . J. Morton, Chem. Comm. (1968) 1565. e n M . Schmeisser a n d K . Brändle, Angew. Chem. 73 (1961) 388. 812 M . Schmeisser a n d E . Schuster, Bromine and its Compounds (ed. Z . E. Jolles), p p . 209-213, Benn, London (1966).

32-95 46-59 3-25

colourless* -175» -45.9a ambientb 4-726» 20-8» + 2-5 + 17-56 b gas solidb

<0*

<0*

38-55 25-31 2-83

gas*

yellow*

yellow* -42d

colourless -107° 18° ambient1* 7.3c 251 c + 6-97 +2208 gasb solidb liquidf

0

b c

ION0 2

BrON02

C10N02

» G. H. Cady, /. Amer. Chem. Soc. 56 (1934) 2635; O. Ruff and W. Kwasnik, Angew. Chem. 48 (1935) 238. R. H. Miller, D. L. Bernitt and I. C. Hisatsune, Spectrochim. Ada, 23A (1967) 223. H. Martin, Angew. Chem. 70 (1958) 97. * M. Schmeisser and K. Brändle, Angew. Chem. 73 (1961) 388. β C. J. Schack, cited in C. J. Schack, K. O. Christe, D. Pilipovich and R. D. Wilson, Inorg. Chem. 10 (1971) 1078. f A. M. Qureshi, J. A. Ripmeester and F. Aubke, Canad. J. Chem. 47 (1969) 4247.

N nmr spectrum Bond dissociation energies (kcal mol_1)b D(0-X) 2>(02N-OX) Force constant,Λ(Ο-Χ) (mdyne A-*)*

14

Colour Melting point (°C) Boiling point (°Q Decomposition temperature (°Q A/fvap(kcalmol-i) Trouton's constant (cal deg"1 mol -1 ) Δ#;, 298[XON02(g)] (kcal mol'i)* AG;, 298[XON02(g)] (kcal ιηοΓΐ)* Infrared spectrum

FON0 2

TABLE 88. HALOGEN(I) NITRATES

1472

CHLORINE, BROMINE, IODINE AND ASTATINE: A. J. DOWNS AND C. J. ADAMS

although the action of C1F on nitric acid produces C10N0 2 in 90% yield813. The nitrate has been used to make other halogen nitrates (Scheme 17). Furthermore, its reaction with metal chlorides (conveniently executed at -78°C in liquid C10N0 2 ) provides a route to anhydrous metal nitrates814. Other reactions are summarized in Scheme 17. Like CIONO2, IONO2 adds across olefinic linkages, and has also been used to oxidize alcohols815. BrR

N205,CFC13 -30°C

Br(NO)

20°C

-—

N02[Br(0N02)2]

BrtX

-20°C

NaNO,

Ti(N03)4 -78°C

02NOCMe2CH,Cl

BrON0 2 + N02[Br(ON02)2] Me 4 N[X(0N0 2 ) n ] (X = Br, n=2; X = I, n=2 or 4)

(NO)2S3Oi0

OCl + N03~

(N0 3 ) 3

I,,CFCL -+- IONO„

<0°C

SCHEME 17. Halogen nitrates

Detailed investigations of its infrared spectrum establish that CIONO2 is structurally akin to FON0 2 . In keeping with the differences in halogen-oxygen and nitrogen-oxygen bond energies, however, the first step in the decomposition of fluorine nitrate is FON0 2 ->F+ON0 2

but chlorine nitrate disintegrates via the reaction C10N0 2 -*C10+N0 2

for which an activation energy of 30 kcal mol _1 has been derived experimentally816. While I(N0 3 ) 3 loses nitrogen dioxide at 0°C, the yellow compound Br(N0 3 ) 3 is more stable, melting with decomposition at 48°C; orange bromyl nitrate, Br0 2 N0 3 , does not persist much above — 78°C. Infrared spectroscopic investigations of the complex nitratohalogen anions [X(N0 3 )J _ (X = Br, n = 2; X = I, n = 2 or 4) could not determine whether the nitrate groups function as mono- or bidentate ligands817. 813 C. J. Schack, Inorg. Chem. 6 (1967) 1938. C. C. Addison and N. Logan, Preparative Inorganic Reactions (ed. W. L. Jolly), Vol. 1, p. 141, Interscience (1964). sis u . E. Diner and J. W. Lown, Chem. Comm. (1970) 333. 816 L. F. R. Cafferata, J. E. Sicre and H. J. Schumacher, Z. physik. Chem., N.F. 29 (1961) 188. en M. Lustig and J. K. Ruff, Inorg. Chem. 5 (1966) 2124. 814

1473

HALOGEN DERIVATIVES OF OXYACIDS

Halogen Perchlorates The interaction of iodine and silver perchlorate in organic solvents (which Gomberg believed to produce CIO4 608) is in fact rather complicated. The various studies of this reac­ tion have been reviewed by Alcock and Waddington609, whose own careful scrutiny led to the findings summarized in Table 89. The course of the reaction depends on the solvent and on the reaction conditions, and in no case can a pure product be isolated. TABLE 89. REACTIONS OF IODINE WITH SILVER PERCHLORATE IN NON-AQUEOUS SOLVENTS

Conditions

Solvent

-85°C

EtOH Et 2 0

EtOH or Et 2 0

Reaction I2+AgC10 4 -> AgI+IC10 4

- 8 5 ° C ; I 2 added to AgC10 4

I 2 +2AgC10 4 -> Agl+Ag[I(C10 4 ) 2 ]

- 8 5 ° C ; AgC10 4 added to I 2

I 2 +AgC10 4 -> AgI+IC10 4 I 2 +2AgC10 4 - * Agl+Ag[I(C10 4 ) 2 ] Ag[I(C10 4 ) 2 ]+IC10 4 -► AgI+I(C10 4 ) 3

room temperature

rapid reaction with solvent

TABLE 90. PROPERTIES OF HALOGEN(I) PERCHLORATES

FOCIO3 Colour Melting point (°C) Boiling point (°C) Decomposition temperature (°C) Infrared spectrum Raman spectrum Force constants (mdyne A~ *) /r(Cl-O) /r(Cl-0) Λ(Ο-Χ) 19 F nmr spectrum

colourlessb -167-3* -15-9/755 mm* ~100c gasc

9-58h

CIOCIO3» pale yellow6 -117±2e 44-5 e ambiente gasf matrix-isolated' solid' liquid' 8-8' 2-65' 2-65'

BrOC103 red« <-78* -20« gasf matrix-isolated'

8-8' 2-65' 2-65'

d

» See also Table 45. G. H. Rohrback and G. H. Cady, / . Amer. Chem. Soc. 69 (1947) 677. c Y. Macheteau and J. Gillardeau, Bull. Soc. chim. France, (1969) 1819. d H. Agahigian, A. P. Gray and G. D. Vickers, Canad. J. Chem. 40 (1962) 157. e C. J. Schack and D. Pilipovich, Inorg. Chem. 9 (1970) 1387. ' K. O. Christe, C. J. Schack and E. C. Curtis, Inorg. Chem. 10 (1971) 1589. e C. J. Schack, K. O. Christe, D. Pilipovich and R. D. Wilson, Inorg. Chem. 10 (1971) 1078. h E. A. Robinson, Canad. J. Chem. 41 (1963) 3021. b

The action of fluorine on concentrated perchloric acid generates the unstable compound fluorine perchlorate818, which has a pronounced tendency to explode, e.g. on freezing. It »is G. H. Rohrback and G. H. Cady, / . Amer. Chem. Soc. 69 (1947) 677; ibid. 70 (1948) 2603.

1474

CHLORINE, BROMINE, IODINE AND ASTATINE: A. J. DOWNS AND C. J. ADAMS

decomposes thermally by two routes: CIF+2O2 * - FOCIO3 -> C 1 0 2 F + 0 2

and readily oxidizes iodide ions: FOCIO3+21- - + I 2 + C I O 4 - + F -

Chlorine(I) perchlorate, which has been briefly treated in connection with the oxides of chlorine (p. 1371), has been used to synthesize bromine599 and iodine perchlorates (Scheme 18). All the halogen perchlorates are hazardous compounds, liable to explode if subjected to thermal or physical shock. BrOSO.F [M

MC1<)4- C's or NO,]

-~-

-20°C

HBr

BrOC103

v C10SO,F

Br +HCK)

-78°C

1 _45°C Br,

or C1F

cs[i(ocio3)4]-

CIOCIO,

-45 °C

1,(033 mole) -50°C

"AgCl Cl2+AgC104

i(ocio3)3

hv

ci 2 ,o 2 ,ci 2 o 7 Cl 2 O 6 (80%),ClO 2 ,Ci 2 ,O 2

SCHEME 18. Halogen perchlorates.

Vibrational spectroscopic studies confirm that the halogen(I) perchlorates have struc­ tures directly related to that of perchloric acid, the hydrogen being replaced by fluorine, chlorine or bromine. While FOCIO3 is apparently the most explosive member of the set, thermal stability decreases in the sequence FOCIO3 > CIOCIO3 > BrOC103 as the polarizability of the terminal halogen increases. Attempts to make I0C10 3 from equimolar amounts of I 2 and CIOCIO3 have been unsuccessful, the iodine being oxidized to the trivalent state819. The reaction of iodine, ozone and anhydrous perchloric acid allegedly affords a material of empirical composition I(ClO4)3,2H2O820. p u r e I(C104)38i9 is a white solid, which is polymeric (according to Raman data) and unstable at temperatures much above -45°C; the mode of decomposition depends on the rate of warming, but the products include chlorine oxides (predominantly C1207), an iodine(V) perchlorate (probably I0 2 C10 4 ) and I 2 0 5 . The pale yellow salt Cs[I(OC103)4] is stable at ambient temperatures: the anion apparently contains a square-planar I 0 4 unit. Halogen Carboxylates With the exception of perfluoroacyl hypofluorites, halogen(I) carboxylates are unstable and cannot be isolated in the pure state. They are intermediates in the Hunsdiecker reaction, 819 K. O. Christe and C. J. Schack, Inorg. Chem. 11 (1972) 1682. 820 F . Fichter and H . Kappeier, Z. anorg. Chem. 91 (1915) 134.

HALOGEN DERIVATIVES OF OXYACIDS

1475

whereby equimolar amounts of the halogen and a silver carboxylate, mixed in an inert solvent, produce an alkyl halide via decarboxylation of the initially formed halogen carboxylate821: _co RC02Ag+X2

> AgX+[RC0 2 X] — » RX (X = Cl, Br or I)

This intermediacy has been verified by the spectroscopic detection of O-Cl bonds in carbon tetrachloride solutions containing CH 3 C0 2 Ag and chlorine822. Aliphatic halogen(I) carboxylates lose carbon dioxide at room temperature, but perfluoro- derivatives undergo decarboxylation only at 100°C. Iodine trifluoroacetate is stabilized as its pyridine complex C5H5N,IOCOCF3, and analogous derivatives of other iodine(I) carboxylates have also been reported; the compounds give only very weakly conducting solutions in acetone. The halogen acetates CIOCOCH3 and BrOCOCH3 are believed to be the effective halogenating agents in mixtures of acetic anhydride with chlorine and bromine respectively823. Oxidation of iodine with fuming nitric acid in the presence of a carboxylic acid or anhydride affords the iodine(III) compounds I(OCOR)3 [R = CHWC13_W, where« = 0, 1, 2 or 3; CF 3 , C 3 F 7 and CÖF 5 ] 8 2 4 . The infrared spectra of the solid perfluoro- compounds are O consistent with the presence of covalent I - O - C ^ units, although in electrolysis of iodine R triacetate the quantity of silver iodide formed at a silvered platinum gauze cathode is in good agreement with Faraday's Law calculations based on the equation I3 + +Ag+3e->AgI

which assumes that P + is present in solution. The iodine(III) carboxylates are quite sensitive to moisture and possess limited thermal stability; thus, the colourless crystals of iodine(III) acetate are fairly stable in the cold, but begin to decompose at 100°C and explode at 140°C825. The compound reacts with methanesulphonic acid to give (CH 3 S0 3 ) 3 I. Other Halogen Oxysalts On oxidation of iodine with concentrated nitric acid in the presence of acetic anhydride and phosphoric acid, the compound IPO4 is obtained. Its nature remains a matter for conjecture, but hydrolysis causes the iodine(III) to disproportionate: 5IP04+9H20 -^ I2+3HI03 + 5H3P04 Moreover, oxyhalogen derivatives of the acids H2SO4, H2S2O7, H2S3OH) and H2Se04 have been prepared by reactions such as and

2C102+3S03 -> [ClO][ClO2][S3O10] HIO3+H2S04 -> I02HS04+H20

The nature of the oxyhalogen unit in these and related compounds has been considered in Section 4A (p. 1352). 821 R. G. Johnson and R. K. Ingham, Chem. Rev. 56 (1956) 219. 822 M. Anbar and I. Dostrovsky, / . Chem. Soc. (1954) 1105. 823 p. B . D. de la Mare, I. C. Hilton and C. A. Vernon, / . Chem. Soc. (1960) 4039; P. B. D . de la Mare and J. L. Maxwell, Chem. and Ind. (1961) 553. 824 M. Schmeisser, K. Dahmen and P. Sartori, Ber. 100 (1967) 1633. 825 N . V. Sidgwick, The Chemical Elements and their Compounds, Vol. II, p. 1244, Clarendon Press, Oxford (1950).

1476

CHLORINE, BROMINE, IODINE AND ASTATINE: A. J. DOWNS AND C. J. ADAMS

C. INTERHALOGENS AND POLYHALIDE ANIONS»26-84i 1. INTRODUCTION

The existence of compounds formed by union of the halogens with one another came to be recognized soon after the identification of the halogens themselves. Thus, the monoand trichloride of iodine were discovered by both Gay Lussac842 and Davy843 in 1813-14. Balard844 first made iodine monobromide in 1826; he also noticed that when bromine is mixed with chlorine the intensity of the colour is much diminished, though a century elapsed before spectroscopic evidence established beyond doubt the existence of bromine monochloride845. First of the halogen fluorides to be identified was iodine pentafluoride, which may have been prepared by Kämmerer in 1862 when he heated dry silver fluoride with iodine84^; however, the true nature of the reaction was not established until 1870: 5AgF+3I 2 -*5AgI+IF 5

and it fell to Moissan first to describe the preparation from its elements and characterization of the pentafluoride846. Working independently, Lebeau and Prideaux discovered bromine trifluoride in 1905 and remarked upon its extreme reactivity847. With the exception of iodine mono- and trifluoride and chlorine pentafluoride, the remaining halogen fluorides were first prepared by Ruff and his collaborators at Breslau between 1925 and 1934: chlorine monofluoride848 and trifluoride849 were made from chlorine and fluorine at 250°C; the 826 j . w . Mellor, A Comprehensive Treatise on Inorganic and Theoretical Chemistry, Vol. II, Longmans, Green a n d C o . , L o n d o n (1922). 827 Supplement to Mellor's Comprehensive Treatise on Inorganic and Theoretical Chemistry, Supplement II, Part I, Longmans, Green a n d C o . , L o n d o n (1956). 828 Gmelins Handbuch der Anorganischen Chemie, 8 Auflage: " C h l o r " , System-nummer 6, Verlag Chemie, Berlin (1927); " B r o m " , System-nummer 7, Verlag Chemie, Berlin (1931); " l o d " , System-nummer 8, Verlag Chemie, Berlin (1933). 829 Gmelins Handbuch der Anorganischen Chemie, 8 Auflage, " C h l o r " , System-nummer 6, Teil B , Lieferung 2, p . 543. Verlag Chemie, Weinheim/Bergstr. (1969). 830 A . G . Sharpe, Quart. Rev. Chem. Soc. 4 (1950) 115. 831 N . N . Greenwood, Rev. Pure Appl. Chem. 1 (1951) 8 4 . 832 R . c . Brasted, Comprehensive Inorganic Chemistry (ed. M . C . Sneed, J. L . M a y n a r d a n d R . C . Brasted), Vol. 3 , p . 179, van N o s t r a n d (1954). 833 H . C . Clark, Chem. Rev. 58 (1958) 869. 834 E . E . Havinga a n d E . H . Wiebenga, Rec. Trav. Chim. 7 8 (1959) 724. 835 E. H . Wiebenga, E . E . Havinga a n d K . H . Boswijk, Adv. Inorg. Chem. Radiochem. 3 (1961) 133. 836 A . G . Sharpe, Non-aqueous Solvent Systems (ed. T . C . Waddington), p . 285, Academic Press (1965). 837 z . E . Jolles (ed.), Bromine and its Compounds, Benn, L o n d o n (1966). 838 L . Stein, Halogen Chemistry (ed. V. G u t m a n n ) , Vol. 1, p . 133. Academic Press (1967). 839 H . Meinert, Z . Chem. 7 (1967) 4 1 . 840 A . I. Popov, Halogen Chemistry (ed. V. G u t m a n n ) , Vol. 1, p . 2 2 5 , Academic Press (1967); MTP International Review of Science: Inorganic Chemistry Series One, Vol. 3 (ed. V. Gutmann), p. 53, Butterworths a n d University Park Press (1972). 841 A . A . Opalovskii, Russ. Chem. Rev. 36 (1967) 711. 842 j . L . G a y Lussac, Ann. Chim. [1] 91 (1814) 5. 843 H . D a v y , Phil. Trans. 104 (1814) 487. 844 A . J . Balard, Ann. Chim. Phys. [2] 3 2 (1826) 337. 845 A . J . Balard, Ann. Chim. Phys. [2] 3 2 (1826) 3 7 1 ; N . V. Sidgwick, The Chemical Elements and their Compounds, p . 1149. Clarendon Press, Oxford (1950). 846 H . Kämmerer, / . prakt. Chem. 85 (1862) 4 5 2 ; G . G o r e , Proc. Roy. Soc. 19 (1870) 2 3 5 ; H . Moissan, Compt. rend. 135 (1902) 563. 847 p . Lebeau, Compt. rend. 141 (1905) 1018; E . B . R . Prideaux, / . Chem. Soc. 89 (1906) 316. 848 o . Ruff, E . Ascher, J . Fischer a n d F . Laass, Z . anorg. Chem. 176 (1928) 258. 849 o . Ruff and H . K r u g , Z . anorg. Chem. 190 (1930) 270.

INTERHALOGENS AND POLYHALIDE ANIONS. INTRODUCTION

1477

850

unstable bromine monofluoride was prepared from the elements at 10°C; and bromine pentafluoride851 and iodine heptafluoride852 were identified as the products of the fluorination of bromine trifluoride (at 200°C) and iodine pentafluoride (at 280°C), respectively. In common with bromine monofluoride, iodine monofluoride cannot be isolated as a pure substance at room temperature since it disproportionates rapidly to iodine and iodine pentafluoride; nevertheless, in 1951 Durie853 identified the IF molecule by its emission spectrum (4350-6900 A) in the bright greenish-yellow flame which accompanies the reaction offluorinewith iodine crystals. As formed by the action offluorineon iodine suspended in trichlorofluoromethane, iodine trifluoride has been represented by Schmeisser and his co-workers854 as a yellow material stable up to — 28 °C, while most recently characterized of the interhalogen compounds is chlorine pentafluoride, first reported by Smith855 in 1963 as the issue of heating to 350°C a 14:1 mixture of fluorine and chlorine at a total pressure of 250 atm. That the donor action of a halide ion on either a halogen or interhalogen molecule gives rise to a polyhalide anion, e.g. I3-, ICl4~ or BrF4~, has likewise been established through a long and sometimes contentious history. Shortly after the discovery of iodine, Pelletier and Caventou856 reported the preparation of a crystalline addition compound of strychnine and iodine, which they called "hydroiodure iodure", and which was un­ doubtedly strychninium tri-iodide, C2iH2202N2,Hl3. The solubility of iodine in potassium iodide solutions and the formation of compounds between metal halides and iodine or an iodine halide attracted early attention to the polyhalides, to which many publications in the nineteenth and early twentieth century bear witness. Cremer and Duncan857 made a careful and thorough re-examination of the methods of preparation, physical properties and reac­ tions in solution and in the solid state of the polyhalides; their systematic survey summarizes most of the early work and adds a considerable weight of new material. In recent years, the range of polyhalide complexes has been substantially enlarged, while a wealth of detail concerning stability and structure has also been brought to light; the outstanding experi­ mental advances have come with the crystallographic analysis of numerous solid complexes, with the development of physicochemical techniques for monitoring reactions in solution, and with the increasing use of non-aqueous solvents as reaction media. It has also become evident that certain halogen and interhalogen molecules are susceptible either to oxidation or to halide ion-abstraction, whereby polyhalogen cations like C12F+, BrF2+ or I 3 + are produced. Such species are the burden of Section 4A and will receive only incidental mention here. Because of their unusual compositions, polyhalogen compounds are not easily described in terms of classical valence rules, and they offer to this day an intriguing test of modern theories of structure and bonding; more than once they have been the subject of theoretical polemics. In this context, there exist numerous analogies between polyhalogen species and, on the one hand, noble-gas halides or, on the other, halide derivatives of the elements wo o . Ruff a n d A . Braida, Z . anorg. Chem. 214 (1933) 8 1 . 851 O . Ruff a n d W . Menzel, Z. anorg, Chem. 202 (1931) 4 9 . 852 o . Ruff a n d R . Keim, Z . anorg. Chem. 193 (1930) 176. 853 R . A . Durie, Proc. Roy. Soc. A207 (1951) 388; Canad. J. Phys. 4 4 (1966) 337. 854 M . Schmeisser a n d E . Scharf, Angew. Chem. 7 2 (1960) 324; M . Schmeisser, W . Ludovici, D . N a u m a n n , P . Sartori a n d E . Scharf, Ber. 101 (1968) 4214. 855 D . F . Smith, Science, 141 (1963) 1039. 856 B . Pelletier a n d J. B . Caventou, Ann. Chim. Phys. [2] 10 (1819) 164. 857 H . W . Cremer a n d D . R . D u n c a n , / . Chem. Soc. (1931) 1857,2243; ibid. (1932) 2 0 3 1 ; ibid. (1933) 181.

1478

CHLORINE, BROMINE, IODINE AND ASTATINE: A. J. DOWNS AND C. J. ADAMS

of Groups V and VI. The striking structural and chemical affinities revealed by the members of formally isoelectronic sequences such as SbX 6 3 ~, TeX62 -, IF 6 ~, XeF 6 ; SbX52 ~, TeX5 -, IF 5 , XeF 5 + ; and SbX4 ~, TeX4 and IF4 + clearly imply that the polyhalogens are representa­ tive of a more general class of molecular aggregate. Typically this class is distinguished (i) by the presence of a central atom coordinated by more ligands than are formally compat­ ible with the achievement of an 8-electron valence-shell, and (ii) by a surplus of valence /?-electrons relative to the number of bonding molecular orbitals which are available through the interaction of valence s- or /^-functions of the central atom. The various theoretical treatments which have been applied to the interpretation of the physical prop­ erties—and notably of the stereochemistries—of polyhalogen species will be dealt with by way of conclusion to the present section. Apart from their intrinsic interest, however, interhalogen compounds are also noteworthy for their action as halogenating agents, a characteristic which has been turned to considerable synthetic advantage, and for the solvent function exhibited by several members, e.g. IC1, BrF3 and IF 5 836 . 2. INTERHALOGEN COMPOUNDS

Introduction In composition, an interhalogen compound stable under conventional conditions belongs to one of the classes XY, XY3, XY5 or XY7; the compounds which have so far been characterized with some degree of assurance are listed in Table 91. To these may TABLE 91. KNOWN INTERHALOGEN COMPOUNDS TypeXY C1F» BrF*>

IF*

BrCl c IC1* IBr»

Type XY 3

Type XY 5

ClF3a BrF3a IFa* I 2 Cl 6 a

ClF5a BrF5a IF5a

Type XY 7 IF7a

a P u r e material long-lived at room temperature; physical properties at least reasonably well defined. b Rapidly disproportionates at room temperature into a mixture of the element and higher fluorides; few reliable physical properties known. c Rapidly dissociates at room temperature into a mixture of the elements; few reliable physical properties known.

be added the recently identified bromine chlorides produced by the action of a micro­ wave discharge on mixtures of chlorine and bromine and trapped in an inert matrix at 20°K858; analysis of the infrared spectra suggests that the trapped molecules include several with T-shaped frameworks possessing either Cl-Br-Cl or Cl-Br-Br "linear" units, e.g. Br2Cl2, BrCl3 and B^Cl. The general stoichiometry of a stable interhalogen compound, that is XYW, is such that n is an odd number, whence it follows that the molecule contains an even number of valence electrons, and the observed diamagnetism is to be anticipated. With the notable exception of the newly discovered Br 2 Cl 2 and Br3Cl858, molecules XY n with n greater than 1 are characterized by a central atom X the atomic 858 L. Y. Nelson and G. C. Pimentel, Inorg. Chem. 7 (1968) 1695.

INTERHALOGEN COMPOUNDS

1479

number of which exceeds that of Y. The greater the atomic number, and hence the size, of X, the more substituents Y it can carry, while the smaller the substituents Y, the more of these can be accommodated around a given X atom: for n to exceed 1, stability appears to demand that Y be either fluorine or chlorine, and for n to exceed 3 that Y be limited to fluorine. Where compounds with higher values of n are known, it is generally the case that the stability is enhanced as the atomic number of X increases. Whether or not a particular interhalogen XY n can be prepared, however, depends on its stability with respect, not only to dissociation into the parent elements X 2 and Y 2 , but also to disproportionation into X 2 and XYW, where m > n. The molecules Br4 and I 4 identified by spectroscopic measurements (see Section 2, p. 1184) formally belong to the interhalogen class XY3; however, in that there are as yet no grounds for believing that the X 4 units embody more than loose association between the X 2 molecules, these appear to have little in common with species like CIF3 or I2C16. More closely related to the interhalogens are derivatives, mostly of the type XY' and XY'3, formed by pseudohalogens Y': examples include XCN, XSCN, XN 3 and I(NCS)3 (X = Cl, Br or I). In many of their properties, e.g. their action as Lewis acids, these pseudohalide compounds resemble conventional interhalogens. Reliable physical properties have been described for but few of the pseudohalides, and, in the absence of sufficient information for a systematic account, such compounds will be represented here, not in detail, but mainly for the comparison they afford with the better defined interhalogen compounds. Radicals of the type XYZ (where X, Y and Z are the same or different halogen atoms), which are short-lived under normal conditions, have been postulated as intermediates in certain gas-phase reactions of halogen and interhalogen molecules. Representative of this class is the radical CI3, which has also been identified in the matrix-isolated condition by its infrared spectrum (see Section 2, p. 1168)859. Molecular-beam studies of atom-recombination reactions also imply the formation of triatomic species of this kind860; according to estimates of dissociation energy and mean lifetime, the most stable configuration of radicals containing different atoms is that with the least electronegative atom in the central position. The thermal or photolytic dissociation of C1F3 is thought to proceed via the formation of C1F2, for which Δ//>° is estimated to be ca. - 1 9 kcal mol - 1 861. Substantiation of this view is, provided by the infrared spectrum reported for C1F2 as produced by the photolysis of chlorine trifluoride or mixtures of chlorine monofluoride and fluorine isolated in an inert matrix at 16°K861; the vibrational characteristics indicate that, whereas CI3 is linear, C1F2 is bent with an apex angle of 136 ± 15°. Such experiments apart, however, interhalogen radicals are the subjects more of speculation than of first-hand evidence. Preparation, Purification and Manipulation 8 ^ -829,838-840,862

All the known interhalogen compounds may be (and most commonly are) obtained by direct combination of the elements. A survey of more specific details for individual compounds, which appears in Table 92, emphasizes that the product formed by the direct 859 L. Y. Nelson and G. C. Pimentel, / . Chem. Phys. 41 (1967) 3671. 860 Y . T . Lee, P. R. LeBreton, J. D . McDonald and D . R. Herschbach, / . Chem. Phys. 51 (1969) 455. 861 J. A . Blauer, H . G. McMath and F. C. Jaye, / . Phys. Chem. 73 (1969) 2683; G. Mamantov, E. J. Vasini, M. C. Moulton, D . G. Vickroy and T. Maekawa, / . Chem. Phys. 54 (1971) 3419. 862 G. Brauer, Handbook of Preparative Inorganic Chemistry; 2nd edn., Vol. 1, Academic Press (1963); Inorganic Syntheses, Vols. 1,3 and 9, McGraw-Hill (1939-67); R. E. D o d d and P. L. Robinson, Experimental Inorganic Chemistry, Elsevier (1954).

I2+C12->2IC1

1.

IF 3 +I 2 ->3IF I 2 +AgF->AgI+IF

I 2 +F 2 ->2IF RI+F 2 ->IF+RF (R = Me or Et)

IC1

or

Br2+Cl2 ^2BrCl

2. 3.

1.

BrCl

IF

BrF3+Br2-»3BrF BrF5+2Br2 -> 5BrF

2.

or

Br 2 +F 2 ->2BrF

1.

BrF

C12+F2->2C1F C1F3+C12-»3C1F

1. 2.

Method

CIF

Compound Copper reaction vessel typically employed. Both methods are effective and useful. CIF is purified by distillation, micro-sublim­ ation or gas chromatography.

Comments

IF detected by its electronic spectrum. It cannot be isolated as a pure compound at room temperature because it disproportionates rapidly: 5IF-*2I 2 +IF 5 However, adducts IF,L (L=2,2'-bipyridyl, quinoline or pyridine) have been reported*.

The reaction may be carried out IC1 is normally prepared by direct combin­ ation of the elements and best purified by with or without a solvent and fractional crystallization of the melt or by at or below room temperature; low-temperature distillation. methods using either liquid or gaseous Cl2 have been des­ cribed.

In the gas phased or in solution, BrCl characterized by spectrophotometric or e.g. in CCl4a-J or waterk, at pressure measurements. The pure com­ normal temperatures. pound cannot be isolated because of the facility of its dissociation into Br2 and Cl2 under normal conditions.

Combustion of I2 or RI with F2 produces a flame; also reaction of I2 with F 2 in CCl3Fat -45°Ce. InCCl3Fat -Ί*°&·Κ -10° to + 30°Ο.

Gas phase. In suitable circum­ BrF detected by its electronic spectrum; furtherfluorinationor disproportionation stances the reaction gives rise of BrF leads to BrF3 or BrF5. to a flame. BrF favoured at elevated tem­ BrF detected spectroscopically but cannot be isolated as a pure compound because of peratures. disproportionation at normal temper­ atures.

Gas phase, 220-250°C. Gas phase, 250-350°C.

Conditions

TABLE 92. THE PREPARATION OF INDIVTOUAL INTERHALOGEN COMPOUNDS

a, c, d, f, 1

a,e,j-l

b,g-i

a,b,f

a-e

References

IF 3

BrF3

2.

1.

2.

1.

I 2 +3XeF 2 -* 2IF 3 +3Xe n

I 2 +3F 2 -»2IF 3

Br2+2C1F3 -> 2BrF3+Cl2 m

Br 2 +3F 2 ->2BrF 3

Copper, nickel, Monel, Inconel or Kel-F apparatus typically employed. Manu­ factured commercially as a technicalgrade chemical. Purification commonly involves treatment with NaF (to remove HF) followed by fractional distillation.

Material purified by fractional crystalliz­ ation of the melt. Dissociation into I 2 and Br2 in the vapour, even at room temperature, makes it difficult to obtain the compound completely free from I 2 and Br2.

A procedure of analytical rather than pre­ parative importance.

Action of F 2 on I 2 suspended in CCl 3 Fat-45°C.

IF 3 is described as a yellow solid stable up to — 28°C; above this temperature it decomposes rapidly to I 2 , IF and IF 5 . Derivatives of IF3are known,e.g. IF3,MFs (M = As or Sb), MIF4 (M = alkali metal or NO) and IF3,L (L = aromatic nitrogen-base).

Reaction between gaseous F 2 "Copper, nickel, Monel or Kel-F apparatus and Br2 vapour, with or with­ employed. BrF3 is manufactured com­ out a diluent, at or near room mercially as a technical-grade chemical. temperature; alternatively, the Purification normally effected by fraction­ F 2 is bubbled into or over al distillation. liquid Br2. C1F3 is passed into liquid Br2 at 0°C.

Gas phase, 20O-300°C.

C12+3F2->2C1F3 C1F+F 2 ->C1F 3

C1F3

or

Direct interaction of the pure elements at or near room temperature.

I 2 +Br 2 -*2IBr

IBr

KIO 3 .

Aqueous acidic solution at room temperature; oxidizing agents which have been used include KMnC>4, chlorine water and

I- + Cl--+ICl+2e

2.

b, h, n, 0

a-d, f, m

a-e

a,c,l

IF 5

BrF5

CIF5

I2CI6

BrCl3, Br2Cl2 and Br3Cl

Compound

Gas phase with excess fluorine, total pressure 250 atm, tem­ perature ~ 350°C. Photochemical reaction carried out at room temperature and 1 atm pressure**. 80-150°C, conversion up to 90% r . 100-300°C, conversion 50-70 %8.

1.

MCl(s)+3F 2 -> MF(s)+C1F 5 (M = alkali metal)

MClF 4 (s)+F 2 -> MF(s)+ClF 5

C1F 3 +F 2 ->C1F 5

hv

C1 2 +5F 2 ^2C1F 5

I2+5F2->2IF5

BrF 3 +F 2 ->BrF 5 KBr(s)+3F 2 -> KF(s)+BrF 5

Br 2 +5F 2 ->2BrF 5

2. Fluorination of I 2 with AgF, CIF3, BrF3 or RuF 5 . 3. HI or metal iodide+F 2 .

1.

2. 3.

1.

4. Electrolytic oxidation of Cl2 or CIF3 in an HF medium*.

3.

2.

or

Compound first prepared by this method.

Reaction carried out in strongly acidic, aqueous solution at <40°C C .

2.1 2 +5C1- + CIO3- + 6H + -> I2C16+ 3H 2 0

I 2 +3C1 2 ->I 2 C1 6

Method (1) is most generally used for the preparation of IF5, though fluorination of I2O5 with CIF3, BrF 3 or SF 4 provides an expedient method of preparation on the small scale. The reaction vessel is typically

Reaction of gaseous F 2 with solid I 2 at room temperature. Various conditions1*. Various conditions1».

Copper, nickel, Monel, Inconel or Kel-F apparatus typically employed. BrFs is manufactured commercially as a tech­ nical-grade chemical. Purification usually involves fractional distillation.

Gas phase · with excess F 2 , temperature > 150°C. Gas phase, temperature 200°C. 25°C, yield - 50% 8 .

Probably the most expedient routes yet devised to C1F5.

A convenient preparative route to C1F5.

Solid I2Cl6 obtained by evaporation of the excess chlorine; cannot be purified by crystallization or vaporization because of its ready dissociation into IC1 and Cl2. Reported as a convenient method of pre­ paring I2C16.

Reaction best carried out bet­ ween I 2 and an excess of liquid Cl 2 at - 8 0 ° C .

1.

Species trapped in solid inert matrices at 20°K.

Action of microwave discharge on gaseous mixtures of Cl2 and Br2.

Br 2 +3C1 2 -> 2BrCl3 etc.

Comments

Method

Conditions

TABLE 92 (com.)

a-d, m, u

a-d, f, m, s

b,e,m,q-t

a,c,f,l

P

References

2.

1.

or

or

I 2 +7F 2 -*2IF 7 IF 5 +F 2 ->IF 7 KI(s)+4F2 -> KF(s)+IF7 MI2(s)+8F2 -> MF2(s)+2IF7 Gas phase, 250-300°O>. Gas phase, 150°CV. 250°CW. e.g. M = Pdx. Method (2) recommended for the preparation of pure IF7 because of the difficulty of drying I2. IF7 reacts with silica, glass, I2Os or traces of water to form OIF5, from which it can be separated only with difficulty. It is normally handled in Monel or nickel apparatus and purified by fractional sublimation.

of copper or nickel; purified IF5 may subsequently be manipulated in Kel-F, silica or (at room temperature) glass apparatus. IF5 is manufactured com­ mercially as a technical-grade chemical. Rigorous purification involves treatment (i) with F2 to oxidize I2 and (ii) with NaFu to remove HF, followed by fractionation . a-c, m, v-x

b c d e f

* Mellor's Comprehensive Treatise on Inorganic and Theoretical Chemistry, Supplement II, Part I, Longmans, Green and Co., London (1956). L. Stein, Halogen Chemistry (ed. V. Gutmann), Vol. 1, p. 133, Academic Press (1967). G. Brauer, Handbook of Preparative Inorganic Chemistry, 2nd edn., Vol. 1, Academic Press (1963). R. E. Dodd and P. L. Robinson, Experimental Inorganic Chemistry, pp. 223-226, Elsevier (1954). Gmelins Handbuch der Anorganischen Chemie, 8 Auflage, "Chlor", System-nummer 6, Teil B, Lieferung 2, Verlag Chemie (1969). Inorganic Syntheses, Vols. 1, 3 and 9, McGraw-Hill (1939-67). * M. Schmeisser, P. Sartori and D. Naumann, Chem. Ber. 103 (1970) 590, 880. h M. Schmeisser and E. Scharf, Angew. Chem. 72 (1960) 324. 1 H. Schmidt and H. Meinert, Angew. Chem. 72 (1960) 109. J N. N. Greenwood, Rev. Pure Appl. Chem. 1 (1951) 84. .k H. Gutmann, M. Lewin and B. Perlmutter-Hayman, / . Phys. Chem. 72 (1968) 3671. 1 Gmelins Handbuch der Anorganischen Chemie, 8 Auflage: "Brom", System-nummer 7, Verlag Chemie (1931); "Iod", System-nummer 8, Verlag Chemie (1933). m H. Meinert, Z. Chem. 7 (1967) 41. n N. Bartlett, personal communication. 0 M. Schmeisser, W. Ludovici, D. Naumann, P. Sartori and E. Scharf, Chem. Ber. 101 (1968) 4214. p L. Y. Nelson and G. C. Pimentel, Inorg. Chem. 7 (1968) 1695. q R. Gatti, R. L. Krieger, J. E. Sicre and H. J. Schumacher, /. Inorg. Nuclear Chem. 28 (1966) 655; R. L. Krieger, R. Gatti and H. J. Schumacher, Z. phys. Chem. 51 (1966) 240. r D. Pilipovich, W. Maya, E. A. Lawton, H. F. Bauer, D. F. Sheehan, N. N. Ogimachi, R. D. Wilson, F. C. Gunderloy, jun., and V. E. Bedwell, Inorg. Chem. 6 (1967) 1918. 8 G. A. Hyde and M. M. Boudakian, Inorg. Chem. 7 (1968) 2648. •E. A. Lawton and H. H. Rogers, U.S. Pat. 3,373,096 (1968). u D. W. Osborne, F. Schreiner and H. Selig, /. Chem. Phys. 54 (1971) 3790. v C. J. Schack, D. Pilipovich, S. N. Cohz and D. F. Sheehan, / . Phys. Chem. 72 (1968) 4697. w H. Selig, C. W. Williams and G. J. Moody, / . Phys. Chem. 71 (1967) 2739; H. H. Ciaassen, E. L. Gasner and H. Selig, / . Chem. Phys. 49 (1968) 1803. * N. Bartlett and L. E. Levchuk, Proc. Chem. Soc. (1963) 342.

IF7

4. I 2 0 5 or metal iodate+F2, C1F3, BrF3 or SF4. Various conditionsb»m.

1484

CHLORINE, BROMINE, IODINE AND ASTATINE: A. J. DOWNS AND C. J. ADAMS

interaction of elemental halogens is acutely dependent on the conditions under which the reaction is carried out. For the formation of certain halogen fluorides, advantage has sometimes been taken of fluorinating agents other than elemental fluorine; possible agents include other halogen fluorides, e.g. C1F3 or BrF3, as well as the noble-gas fluorides KrF 2 and XeF 2 838 , platinum and plutonium hexafluorides838»863, and dioxygen difluoride838. Another variant of the normal convention involves fluorination of a suitable metal halide, usually by elemental fluorine; chlorine and bromine pentafluorides and iodine heptafluoride are expediently produced by methods of this sort. Thus, for the preparation of pure iodine heptafluoride, fluorination of potassium iodide has been recommended864 in pref­ erence to the reaction between iodine and excess fluorine, primarily because iodine is difficult to dry, with the result that the product of the latter reaction tends to be contaminated with iodine oxypentafluoride, OIF 5 . Reactions of the type X2+XF3^3XF

have also been exploited in the formation of the halogen monofluorides. Chlorine monofluoride, free from the trifluoride, is thereby efficiently produced at 250°C, and there are good reasons for believing that the monofluorides of bromine and iodine are formed in a similar fashion, albeit fleetingly under normal conditions. Without exception, the interhalogen compounds are highly reactive materials, seldom easy to purify or to manipulate in the pure condition. The destructive action on organic materials also makes the compounds highly injurious to living matter. The toxicity of the vapours depends upon their virulent action on the eyes, nose and throat, the bronchi and mucous membrane being especially vulnerable to attack; the liquids likewise have a drastic physiological effect on the skin, giving rise to painful burns and necrosis. Precise details of the biological effects are sparse, though qualitative reports of the action of certain halogen fluorides have been published838»862»865, and procedures have been described for the safe handling and monitoring of toxic levels of chlorine trifluoride838»866. With respect to dissociation, the compounds BrCl, IC1, IBr and I2C16 are of relatively low stability at normal temperatures. Such dissociation renders impossible the production of bromine monochloride free from elemental chlorine and bromine, and inhibits the use of distillation for the purification of the other compounds. Following their preparation, iodine monochloride and monobromide are commonly refined by fractional crystallization of the melt827»862, but the dissociation pressure of iodine trichloride is such as to frustrate even crystallization as a method of purification. Accordingly, iodine trichloride is normally produced under conditions which ensure that contamination is restricted to materials like chlorine, which can be removed simply by vaporization827»862. All of these interhalogen compounds attack not only metals, but also cork, rubber and many other organic materials. They are best handled in vacuo or in an inert atmosphere using glass apparatus which excludes grease or moisture, storage being accomplished preferably in sealed glass ampoules. By contrast, the fluorides BrF, IF and IF 3 are susceptible to disproportionation which is rapid at room temperature; on the evidence presently available, it is doubtful whether 863 F . P. Gortsema and R. H. Toeniskoetter, Inorg. Chem. 5 (1966) 1925. 864 H . Selig, C. W. Williams and G. J. Moody, / . Phys. Chem. 71 (1967) 2739. 865 H . C. Hodge and F. A. Smith, Fluorine Chemistry (ed. J. H. Simons), Vol. IV. Academic Press, New York and London (1965). 866 L . M. Vincent and J. Gillardeau, Comm. Energ. At. (France), Rapt. CEA No. 2360 (1963).

INTERHALOGEN COMPOUNDS

1485

pure samples of any of these compounds have yet been prepared. Although the welldefined fluorides CIF, C1F3, C1F5, BrF3, BrF5, IF 5 and IF 7 are prone neither to dissociate nor to disproportionate at normal temperatures, their susceptibility to reaction with moisture, glass, organic materials and most metals is such that their preparation, manipula­ tion and storage calls for much the same stringency of operation that is demanded by fluorides like XeF2, ReF7 and PtF6 867,868. According to numerous reports, several of the halogen fluorides, e.g. C1F3, BrF3 and IF7, attack Pyrex glass or quartz even at room temperature, while this becomes the general pattern of behaviour at elevated temperatures. It is probable, however, that impurities like hydrogen fluoride initiate or accelerate this reaction, as it has also been reported, for example, that very pure chlorine trifluoride has no effect on Pyrex or quartz at normal temperatures838. Materials found to offer the best resistance to attack include Monel, Inconel, nickel, copper, stainless steel and sapphire, together with Kel-F and Teflon plastics, which, despite their instability at elevated tempera­ tures, are frequently used for the fabrication of gaskets, valve-seatings and both reaction and storage vessels. The handling of materials like BrF3 and IF 7 requires, ideally, a vacuumsystem having a working manifold, valves and a Bourdon or bellows manometer constructed in Monel, nickel, copper or stainless steel; access is thence gained to traps or reaction vessels made either in metal or Kel-F, depending on the severity of treatment to be applied to the fluoride. For investigations of the infrared absorption of a halogenfluoride,a cell having a nickel body and bearing, typically, AgCl, CaF2 or polythene windows has been effectively employed868, while, to examine the Raman spectrum and certain other properties of the vapour or liquid, silica vessels have been extensively used, though baking in vacuo and preliminary exposure to thefluorideare recommended for the pre-seasoning of the equipment. The electrical conductivity of the pure liquid or of solutions of other materials in the liquid is probably best measured with the aid of a conductivity cell made in Kel-F869, though earlier investigations relied, with varying degrees of success, on glass or silica cells. Impurities likely to be contained in freshly prepared or commercial samples of the halogenfluoridesinclude the elemental halogens, hydrogenfluoride,other halogen fluorides and oxyhalogen derivatives, e.g. OIF5. To purify a particular halogen fluoride, the two measures most commonly taken are (i) treatment with an anhydrous alkali-metal fluoride to remove hydrogen fluoride and certain of the halogen fluoride impurities in the form of involatile salts, e.g. NaHF2, and (ii) fractional distillation or sublimation, either from trap to trap or via a suitable low-temperature column. For example, the samples of iodine pentafluoride used for recent calorimetric and vapour-pressure measurements were derived from the commercial product in the following stages: first, treatment at room temperature, in a Kel-F container, with a small amount of fluorine to oxidize elemental iodine to the pentafluoride; second, removal of traces of hydrogen fluoride by heating to 150°C in the presence of finely divided and thoroughly dried sodium fluoride; and finally, fractionation of the volatile material to separate the pentafluoride from the more volatile heptafluoride870. Similarly, for the most recent vapour-pressure measurements on iodine heptafluoride871, 867 D . F. Shriver, The Manipulation of Air-sensitive Compounds, p. 105, McGraw-Hill (1969); B. Weinstock, Record of Chemical Progress, 23 (1962) 23; H. H. Hyman (ed.), Noble-gas Compounds, The University of Chicago Press, Chicago (1963). 868 J. H. Canterford and T. A. O'Donnell, Technique of Inorganic Chemistry (ed. H. B. Jonassen and A. Weissberger), Vol. 7, p. 273. Interscience, New York (1968). 369 L. A. Quarterman, H. H. Hyman and J. J. Katz, / . Phys. Chem. 61 (1957) 912. S70 D. W. Osborne, F. Schreiner and H. Selig, / . Chem. Phys. 54 (1971) 3790. 871C. J. Schack, D . Pilipovich, S. N. Cohz and D . F. Sheehan, / . Phys. Chem. 72 (1968) 4697.

1486

CHLORINE, BROMINE, IODINE AND ASTATINE! A. J. DOWNS AND C. J. ADAMS

the sample was prepared by the fluorination of iodine pentafluoride and purified (i) by treatment with anhydrous potassium fluoride to abstract both hydrogen fluoride and iodine pentafluoride (as KHF2 and KIF 6 , respectively), and (ii) by fractional sublimation in vacuo. The use of gas chromatography has also been reported872 for the separation of mixtures of halogen fluorides, fluorine and other halogens, hydrogen fluoride and other volatile fluorides (e.g. UF^); a typical assembly for this purpose incorporates a nickel column which is packed with Kel-F powder coated with Kel-F oils. Pseudohalide analogues of the interhalogen compounds are commonly prepared by the action of an elemental halogen or an interhalogen compound on a derivative of the pseudohalide anion827»832»862»873'874, e.g. M C N + X 2 -+ XCN + MX (M = alkali metal or £Hg; X = Cl, Br or I; aqueous or non-aqueous media) MN3+X2-*XN3+MX (M = alkali metal, Ag or £Hg 2 ; X = Cl, Br or I; various conditions) 3IC1+3MECN -> 3MC1+I(NCE) 3 +1 2 (M = K, NH 4 , Ag or £Pb; E = O or S)

The products range from the toxic but comparatively stable cyanides XCN to the highly explosive azides XN 3 . A tendency to suffer polymerization appears to be a common characteristic, e.g. 3XCN->[XCN] 3 Halogen Cyanuric cyanide trihalide (X = Cl or Br) 2XNCO -> X 2 N C O NCO

Physical Properties 8 2 7 » 8 2 9 » 8 3 1 » 8 3 5 ' 8 3 8 " 8 4 0

1. General characteristics. In listing the principal physical properties of the interhalogen compounds, Table 93 alludes to structural, thermodynamic and spectroscopic parameters of the molecules, as well as to characteristics of the condensed phases, notably melting and boiling points, vapour pressure, density, viscosity and electrical conductivity. However, it should be emphasized that the data are not all equally reliable. In some instances, e.g. that of bromine monochloride, thermal dissociation precludes the possibility of accurate measurements of properties like the melting and boiling points, and the figures quoted merely represent a very approximate estimate of the range of existence of the liquid phase. The facility of disproportionation of the monofluorides BrF and IF likewise frustrates measurements of meaningful physical properties for the condensed phases, though the molecules have been characterized spectroscopically in the vapour phase. In the light of this, the significance of the physical properties originally reported for bromine monofluoride by Ruff and Braida850 must be regarded as dubious. For iodine trifluoride, numerical physical properties have yet to be described. 872 j . F. Ellis, C. W. Forrest and P. L. Allen, Anal. Chim. Acta, 22 (1960) 27; J. G. Million, C. W. Webber and P. R. Kuehn, U.S. Atomic Energy Commission Rept. K-J639 (1966). 873 K . Dehnicke, Angew. Chem., Internat. Edn. 6 (1967) 240. 874 A . Hassner, M. E. Lorber and C. Heathcock, J. Org. Chem. 32 (1967) 5 4 0 ; F. W. Fowler, A . Hassner and L. A. L e w , J. Amer. Chem. Soc. 89 (1967) 2077; A . Hassner and F. Boerwinkle, ibid. 9 0 (1968) 216.

Ground state properties Internuclear distance, r e (A)n Dipole moment, D Vibrational frequency, o>e(cm_1) Anharmonic vibrational constant, avc^cm" 1 ) Force constant, ke (mdyne/A) Dissociation energy: A>° Thermodynamic properties Δ#,°[ΧΥ( β )] at 298°K (kcalmol - 1 ) AG,°PCY(g)] at 298°K (kcalmol" 1 ) S°[XY(g)] at 298°K (caldeg^mol" 1 ) Δ#,°[ΧΥ( δ )] at 298°K (kcalmol" 1 )

54-451

Molecular weight10 Gaseous molecules Electronic ground state configuration cßir+n**, *ΐΣ+ Electronic transitions, r e (cm _ 1 ), relative to the 2Τ*Σ+ state (excited state in parentheses)11

—

—

54-708

-17-68

-13-88

52068

-1408

2-80 kcal eV 51-4Ϊ3 2-2313

3-600

4089 kcal eV 59-428 2-5778

4-484 kcal eV 60-358 2-6178

-13-58

1-611

2·4 8 π

4·5ΐι

—

—

57-36"

-0-23"

-28-18 56-458

+ 3-50Ϊ4

eV ?2·878

-22-68

kcal ?66·28

440"

6-23 (35C1F)11

6Ο8Ί911

672-611

786-34 (35C1F)11

2-138 0-5711

1-756 1-2911 1·9089 —

16,795 (3Π 0 +)

19,053 ·7 5 (3Π 0 +)

18,281-2 (3Π 0 +) 17,385 (3Πι)

115-357

BrCl

61,563 (D) 59,318 (C)

145-9029

IF

61,615(11) 57,900 (I)

98-902

BrF

1-6281 0-88111

18,956 (3Π 0 +)

C1F

Property

(a) Type XY

TABLE 93. PHYSICAL PROPERTIES OF THE INTERHALOGENS1"9

m

eV 1-81713

2071

61-822" -2-5"

-8-44(a)is.i6

+0-89"

+ 9-76"

kcal 41-9P3

0-83 (Ρ9ΒΓ)11

268-71 (FöBr)11

2-485 l-21 12a

56,369* (G) 51,701* (F) 39,126 (E) 38,713 (D) 35,427 (C) 16,814 [B'(0 + )] 16,165 (3Π 0 + ) 12,213 (3Π0

206-8085

IBr

59140"

—1-3715.16

+4-18«.™

2-386 kcal eV 49-63Ϊ3 2-152Ϊ3

1-501 (I35C1)11

384-293 (P5C1)11

2-320912b l-24 12b

~ 18,000* [B'(0 + )] -17,344* (3Π 0 +) 13,556-21 (3Πι)

58,168* (D) 53,457* (C) 37,741 (E)

162-3575

IC1

C1F

+

+

4-80 (173-0°K)8

Heat of vaporization (kcalmol - 1 )

-155-68

°C -10018

117-58

°K 17308

—

Melting point

ΧΥ (2Σ )^-ΧΥ(ΐΣ ) Properties of the condensed phases Boiling point

+

Thermodynamic properties (cont.) AGf°\XY(s)] at 298°K (kcalmol"i) — S°pCY(s)] at 298°K (caldeg"imol~i) — AHf°\XY(\)) at 298°K (kcalmol"i) — AG,°[XY(1)] at 298°K (kcalmol' 1 ) — 5°[XY(1)] at 298°K 1 (caldeg^imol" ) — Heat capacity, Cp° -1 (cal deg mol~i) 6-646-9-722 Gas (100-6000°K)i7 Liquid (303-68-317-76°K)i5 — Solid (17·71-295·04°Κ)ΐ5 — (263-273°K)5 — lonization potentials: kcal eV + 1 + ΧΥ (2Π 3 / 2 )^-ΧΥ( Σ ) ^29319 12-719 + + ΧΥ (2Π 1 / 2 )^-ΧΥ(ΐΣ ) ΧΥ + (2Π)<-ΧΥ(ΐΣ + ) —

Property

(a) Type XY

TABLE 93 (cont.)

— 6-983-10-146

— — —

— 6-970-9-704

— — —

—

—

11-819

°K °C ~ 2932 ~ 202 Disproportionates - 2 4 0 2 - -332

27219

—

—

10-519

°K °C Disproportionates

24219

eV

—

—

kcal

—

—

eV

—

—

kcal

—

IF

—

BrF

eV

—

—

11-119

320-518

13-9018

1-91-13-74(a) ca. 13-7605) kcal eV 226018 9-8018 241018 10-4518 284-818 12-3518

24-63-24-61

7-213-9-580

13-9903) 9-876 (liquid at 298·15°Κ)ΐ5.ΐ6 12-622 (solid, α at 298·15°Κ)ΐ5.ΐ6

287-09

°Κ °C ~ 3899 ~ 1169 Dissociates 3149 419

kcal eV 220-918 9-5818 235-718 10-2218 265-218 11-50 ( 2 Π 3/2 ) 18 309018 13-4018

— —

—

7-623-9-620

—

— 32-314

—

-3.314i5.i6

—

—

IBr

-5-6915.16

23·405(α)ΐ5.ΐ6

-3·334(α)ΐ5.ΐ6

IC1

°K °C °K °C ~ 2782 ~ 52 370-3737 97-1007 Dissociates Some dissociation -2072 - - 6 6 2 300-5315 27·38ΐ5(α)

25619

kcal

— —

—

7-101-9-540

—

—

—

—

—

BrCl

logp = 15-7383109/Γ+ 1-538Χ105/Γ2 (liquid, 1 2 3 168°K)8

—

280

— —

— —

— —

12-60 4-38 11-90 119-5° 2-37 2-44 3 00 3 08

t

8-883 8-400 7-568 91-35° 2-35 2-44 2-94 3 06

ß

Partial pressure IC1(1): 38-52135-5(303-60330°K)i6 ICl(s):0-4432-62(250300-53°K)i6 Monoclinic, space group P2i/c (a20 and 02*) 8 (<* and β)

-26-5 2·773(α)ΐ5 2-27(0)9

tt

3-18 3-76

2-52

4-903 6-993 8-931

4

Orthorhombic, space group Ccm2i(no. 36)22

— —

f Structures consist of zigzag chains composed of two types of ICl molecule: (i) with both I and Cl atoms engaged in short-range intermolecular interactions along the chain, and (ii) with the Cl atom branching off the chain and not involved in such interactions (Fig. 36). The closest approach between atoms in different molecular chains implies normal van der Waals' interaction, cc-form: chains puckered and branching Cl atoms eis to each other2o. ß-form: chains essentially planar and branching Cl atoms trans to each other2i. 11 Molecules form a herring-bone pattern similar to that in crystalline l2 22 .

Intermolecular distance (Ä)

Intramolecular distance (Ä)

ß

a b c

Molecules per unit cell Unit cell dimensions (A)

Crystal structure

Vapour pressure, p(mm Hg)

Entropy of vaporization at boiling point (Trouton constant) (cal deg"1 mol" 1 ) Heat of fusion (kcal mol ~ i)

Magnetic susceptibility, X( X106 c g s units)

Viscosity of liquid (centipoise) Electrical and magnetic properties Specific conductivity of liquid (ohm - 1 cm _ i)

Liquid

Mechanical properties Density (gem"3): Solid

Solid Asymmetry parameter for the solid, η 1291 Mössbauer spectrum: isomer shift relative to standard ZnTe source (cm sec - 1 ) Nmr spectra

Spectroscopic properties Vibrational spectra of condensed phases Quadrupole coupling constants, e 2 ö#(MHz): Gaseous molecules24

Property

(a) Type XY

TABLE 93 (cont.)

— —

—

—

—

—

1-9x10-7 (145°K)8

—

—

—

—

—

—

1-62(173°K)2

—

19F resonance studied8

—

— —

—

—

—

—

—

I, -3438-15

—

127

79ßr, 10890

35C1, - 1 4 6 0

—

IF

—

BrF

Refs. 8, 23

C1F

—

—

—

—

— —

—

—

35C1, -103-6 79ßr, 876-8

Ref. 23

BrCl

—

- 5 4 - 5 (s, 17°C)27

4-403-1-817x10-3 (26-75-122-0°C)2

4-19 (28°C)7

—

302-40x10-4 (40-55°C)7.9

—

4-4157-4-4135 (273-283°K)5 3-7616-3-7343 (315-323°K)5

012326

—

-00626

79ßr, 722 1271, - 2 7 3 1 1271, -289226

Ref. 23

IBr

017326

0029825

3-85(a,0°C)9 3-66(j3,0°C)9 3-18223/[l + 91590x10-4/+ 8-3296x10-7/2+ 2-7501x10-9/3], / in °C5

127

C1, -85-84 1 2 b I, -2927-87 12b i27i,_ 3046 (77°K)25

35

Ref. 23

IC1

TABLE 93 (cont.)

Y

Θ^Ικ

Y

Ionization potential

19

Thermodynamic properties: AH,° at 298°K (kcal mol" 1 ) AGf° at 298°K (kcal mol" 1 ) S° at 298°K (cal deg"1 mol" 1 ) Mean X-Y bond energy of XY3 molecule at 298°K (kcal) Heat capacity of gas, Cp° (cal deg-i mol-i) (100-6000°K)i7

Λ Λ

V6 Φ2) Valence stretching force constants (mdyne/A)28»29:

v4(W v5(W

Dimensions: r{k) R{k) Θ Dipole moment, D Vibrational frequencies (cm""1): n ifli) nifli) *3 ifli)

Y

Molecular weight™ Ground state properties of gaseous XY3 molecule Structure: T-shaped molecule, Civ symmetry

Property

(b)TypeXY3

kcal 297

eV 12-9

9-347-19-859

9-394-19-857 eV 13 0

48-2

41-7

kcal 300

3 009 4-084 BrF3(g) -61-09" - 54-84" 69-89"

2-704 4193 C/F3(g) [C/F3]2(g) -39.3530 -81-11« -29-7530 -56-7" 67-28" 117"

1-8108 1-7218 86°12-6,8 1198 Refs. 28, 29 675 552 242 612 350 242

136-899

92-448

1-6988 1-5988 87°29'8 0-5578 Refs. 28,29 752-1 529-3 328 702 442 328

BrF 3

CIF3

kcal 265

eV 11-5

—

— —

— —

—

— -66*

—

— — — — — —

— — — — — —

/*3 116* 110* -72*

—

466-527

I2C16

— — —

183-8995

IF 3

Molecular geometry and dimensions

Molecules per unit cell Unit cell dimensions

Crystal structure

Thermodynamic properties: Ulf at 298°K (kcal mol" 1 ) AG,° at 298°K (kcal mol" 1 ) 5° at 298°K (cal deg"1 mol" 1 ) Heat capacity of solid, Cp° (cal deg"1 mol" 1 )

Boiling point Melting point Transition temperature (solid I to II) Heat of vaporization at boiling point, A/fvap (kcal mol" 1 ) Entropy of vaporization at boiling point (Trouton constant) (cal deg"1 mol" 1 ) Heat of fusion (kcal mol" 1 ) Vapour pressure of liquid, p(wm Hg)

Properties of the condensed phases

Property

(b) Type XY3

TABLE 93 (com.)

Planar I2CI6 molecules, point group Dm ^(I-Clterminal) = 2-38,2-39Ä rf(I-Clbrld8e)=2-68,2-72Ä Z_ ^terminal - ! - ^terminal = 94° Z.Clbridge""I~Clbridse β 8 4

ZF a x -Br-F e q = 8 2 0 , 88-4°

^ F a x - C l - F e q = 86°59'

Triclinic, space group ΡΪ34 1 I2CI6 molecule a = 5-71 A a = 130°50' b= 10-88 A 0 = 8O°51' c = 5-48Ä y=108°30'

1-90-24-50

(IS-^OOK)1*

T-shaped BrF3 molecules, point group Civ
Orthorhombic, space group Cmc2x (148°K)33 4 a = 5-34 A b = 7-35 A c = 6-61 A

—

T-shaped CIF3 molecules, point group C2»
Orthorhombic, space group P/ima (153°K)32 4 a = 8-82 5 A £ = 609 A c = 4-52Ä

—

—

—

— —

/C73(s) -21-340 1 6 -5.41316 40-3951*

— — —

log/? = 7-748531685-8/(f+220-57) t = 38-72 to 154-82°C8 BrF3Q) -71-914 -57-5" 42-6"

23108 1-81938

logp = 7-367111096·917/(ί+232·75) / = -46-97 to +29-55°C8 C/F3(l) -45-6530

10-2358

6-5808

— —

—

— —

-28°C3i

°C Dissociates 3849 101 (16 atm)9

°K

I2CI6

25-78 2-8758

—

—

Described as a yellow solid which decomposes at

IF 3

—

-82-668

190498

°C 125-758 8-778

°K 398-908 281-928

BrF3

—

°C 11-758 -76-328

°K 284-908 196-838

C1F3

Surface tension of liquid (dyne cm" 1 )

i^F chemical shift of gas (ppm rel. S1F4) /(F-F)(Hz) Mechanical properties Density, d(g cm" 3 ): Solid Liquid

Asymmetry parameter, η 1291 Mössbauer spectrum of the solid: isomer shift relative to standard ZnTe source (cm sec - 1 ) i^Fnmr spectrum: Chemical shift of liquid (ppm rel. CF3COOH as an external standard) /(F-F)(Hz)

Spectroscopic properties Vibrational spectra of the condensed phases Halogen nqr spectra of the solid: e2Qq(UHz)

26-6-18-7(273-2323-l°K)39

2-53 (153°K)32 d4l = 1-8853-2-942 x 10"3/-3-79x 10-6/2 (/= -5°to46°Q8

-313-24, -187-32 (300°K)37 441 (300°K)37

- 1 9 2 - 2 , - 8 1 0 (210°K)37 422-2 (210°K)37 At higher temperatures the multiplets coalesce ultim­ ately giving a single broad resonance

35C1,150-2624

Refs. 8, 28 and 29

37-1-33-8(285-2318-2°K)40

3-23 (292°K)8 2-8030 (298°K)38 2-7351 (323°K)38

Single resonance at normal temperatures

-54-3(~300°K)38

Refs. 8, 28 and 29

Ref. 35a

0-35026

dA-*> = 3-203; i/ 4 15 =3-11075

35

Clterminal, -68-3436 Cl brldjre , +27-4836 !27if 3035 (77°K)25 !27I, 0077225

35

Ref. 35b

Magnetic susceptibility, X ( x 106 c g s units) Molar refraction (cm* mol~i)

Liquid

Specific conductivity (ohm - 1 c m - 1 ) : Solid

Gas

Electrical, magnetic and optical properties Dielectric constant,«: Liquid

Mechanical properties (cont.) Viscosity of liquid (centipoise)

Property

(b) Type XY3

TABLE 93 (cont.)

- 2 6 - 5 (1, 300°K)44 10-34 (g, 299°K)8

4-9 x 10-9 (298°K)8

—

= 4-754-0-018/ 1 (t = 0-42°Q4i 1 002825-1 001929 (319-3-413-3°K)42

€t

0-448-0-316 (290-5323-2°K)39

C1F3

- 90-2 (ICI3 solid, 290°K)27

— —

8-34x10-3-7-31x10-3 (283-7-323-2°K)2.43 - 3 3 - 9 ( 1 , 300°K)44 12-92 (g, 326°K)8 13-22 (1,298°K)8

0-019x10-3-8-35x10-3 (343-373°K)2 8-60x10-3-0-36x10-3 (375-420°K)2

—

—

—

1 003748-1 002964 (415-5-448-2°K)42

—

—

—

I2C16

—

—

IF 3

—

3-036-1-775(286-4312-8°K)40

BrF3

Vibrational frequencies (cm - 1 )

Dipole moment, D

Molecular weight™ Ground state properties of gaseous XYn molecule Structure

Property

(c) Types XY5 and XY7

TABLE 93 (cont.)

Square pyramid, CAV symmetry45 ί/(ΒΓ-Ρ ΑΡΐοα1 )=1·689Α
Square pyramid, C*v symmetry8
Refs.48,49 *i(ai) = 709(R,i.r.) v 2 (ai) = 539(R,i.r.) v 3 t o ) = 495(R,i.r.) v 4 ( W = 480(R)i v 5 ( W = 346*(R) v 6 ( « = 375(R)i vn{e) = 7 2 5 ( R , i . r . ) vs(e) =484(R,i.r.) =299(R,i.r.) V9(e) 1 Refers to liquid-phase. Other frequencies refer to gas-phase

174-896

130-445

Ref. 48 vx (αι) = 683 (R, i.r.) v 2 (ai) = 587(R,i.r.) v3 («1) = 369 (R, i.r.) y 4 ( W = 535(R)i v5tf>i) = 281*(R) v6(^2)=312(R)i =644(R,i.r.) Vl(e) =415(R,i.r.) vs(e) v 9 (e) = 2 3 7 ( R , i . r . ) i 1 Refers to liquid-phase. Other frequencies refer to gas-phase

1-51»

BrF5

C1F5

* 3 (*ι) = 318(R,i.r.) v4(b1) = 602(R) v 5 tf>i)=181*(R) v 6 (6 2 ) = 274(R) v7(e) = 631,635 (R, i.r.) v8(
V 2 (ÖI) = 614(R,i.r.)

Refs. 50, 51 "1 (*i) = 710 (R, i.r.)

2-188

Square pyramid, Qv symmetry45
221-8964

IF 5

Some evidence of polar conformations at low temperatures Ref. 52 νι (an = 676 (R) v 2 (ai') = 635(R) *3 ißn = 670 (i.r.) v 4 (*2")=365(i.r.) v 5 (*i') = 746(i.r.) v 6 (e!') = 425(i.r.) v 7 ( ^ ' ) = 257(i.r.) v8 (*!*) = 510 (R) v 9 (eiO = 352(R) "io(*i") = 310(R) v u (*2*) inactive All refer to gas-phase

04?

Pentagonal bipyramid approximating to Dsh symmetry46 ^ a - F a « ) = 1-786 A
259-8933

IF 7

Solid

Melting point Triple point Transition temperature (solid I to II) Heat of vaporization at boiling point, Δ// ν α ρ (kcal mol" 1 ) Entropy of vaporization at boiling point (Trouton constant) (cal deg - 1 mol - 1 ) Heat of fusion (kcal mol" 1 ) Vapour pressure, /?(mm Hg): Liquid

Ionization potential19 Properties of the condensed phases Boiling point

Valence stretching force constants (mdyne/A): Apical X-F bonds Basal or equatorial X-F bonds Thermodynamic properties of gaseous molecules: AHf° at 298°K (kcal mol" 1 ) AGy° at 298°K (kcal mol" 1 ) S° at 298°K (cal deg"1 mol" 1 ) Mean X-F bond energy of XF n molecule at 298°K (kcal) Heat capacity of gas, Cp° (cal deg"1 mol - 1 )

Property

(c) Types XY5 and XY7

TABLE 93 (cont.)

22-7651 2-6825!

log/? = 7-4837-1197/Γ (192-95-391-05°K)53

—

9-42151 9-40351

log/? = 6-4545+0001101/ log/> = 2 9 0 2 1 6 7 -895/(/+206)8 (tin °C) 309014/Γ-6-96834 log T^ 01-1-7(185-208°K)2 log/> = 11-764-3035/Γ8

—

23-38

—

20-453

—

282-57151 282-55351

8-59551

—

—

-60-58

°C 41-308

7-318

—

—

212-658

—

11-830-31-767 (lOO-OOOO^K)1? 311 kcal (13-5 eV) °K °C 377-6351 104-4851

—

—

~-1208

6-458

log/7 = 7-49671291-58/Γ56 log/> = 11-23193046·93/Γ+197,769/Γ2 (195-273°K)57

—

21-156

5-9156 A#subl = 6057

-1538

279-608

—

°K °C 4-778 277-928 (sublimation temperature)

13-213-43-682 (10Ο-60Ο0°Κ)ι:7

55-4

-229-98.51 -201-38.51 87-408

-200-85 1 -184-6 5 1 80-4551 64-2

4·46 52 3·7852

IF 7

4-8249 3-8249

IF 5

5-31353

—

—

—

—

- 1 0 3 ±455

170±455

°K 314-458

10-839-31-770 (ΙΟΟ-όΟΟΟ'ΊΟ1?

20-92-30-85 (250-1000°K)54 °C -13153

44-7

36-9

—

-102-5i4 - 8 3 -8" 76-50H

-60-9+4-553 -390±4-553.54 74-354

°K 2600553

4-0349 3-2449

BrF5

3-4749 2-6749

C1F5

^1

Molecular geometry and dimensions

Molecules per unit cell Unit cell dimensions

Crystal structure

Solid

Critical properties Thermodynamic properties: AHf° at 298°K (kcal πιοΓΐ) Δ(?,° at 298°K (kcal mol'i) S° at 298°K (cal deg"i mol'i) Heat capacity, Cp° (cal deg"1 mol-i): Liquid Ref. 53

=

20 1507 A 6-836 A 18-24 A 92-96°

Z.Faploar"Br-Fbaeal 80-5-86-5°

A similar unit is also found in crystalline XeF2,IF559.

L Faplcal-I-Fbasal) = 81-9°

a = b = c = ß=

41-92-41-55 (282-571-350°K)5i 0-243-32-75 (6-282-571°K)5i Monoclinic, space group most probably c2/c59

IF5(\) -210-851 -186-651 53-7451

Square-pyramidal IF5 Square-pyramidal BrFs molecules: molecules:
a = 6-422 A b = 7-245 A c = 7-846 A

Orthorhombic, space group Cmc2i33

BrFs(\) -109-614 -84-1" 53-814

The data are compatible with the presence of IF7 pentagonal bipyramids with
&?//
Gas

Electrical, magnetic and optical properties Dielectric constant,«: Liquid

Surface tension of liquid (dyne cm "i) Viscosity of liquid (centipoise)

Liquid

Solid (Mössbauer spectrum)6i 1291 Mössbauer spectrum of the solid: isomer shift relative to standard ZnTe source (cm sec"1) i^F nmr spectrum: Chemical shift of liquid (ppm) Basal or equatorial F atoms Apical F atoms /(F-F)(Hz) Mechanical properties Density, d(g cm ~ 3): Solid

Spectroscopic properties Nuclear quadrupole coupling constants, e2Qq(MHz): Gas (microwave spectrum)6**

Property

(c) Types XY5 and XY7

TABLE 93 (cont.)

c = 308-0015/ (/ = - 8 0 ° t o - 1 7 ° C ) 5 3

— —

3-553-1-396Χ10-2Γ + 4-565 Χ10-5Γ2 -6-311 X10" 8Γ3 (193-372°K)53

1-00279 (297-05°K)53

d=

—

-247relCFCl 3 5 5 -412relCFCl 3 5 5 13055

—

—

—

C1F5

2-5509-3-484x10-3/ -3-45x10-6/2 ( / = -15°to76°C)8

1006320-1004378 (345-6-430-8°K)64

e = 8-20-0-0117/ ( / = -ll-7°to24-5°C)64

24-3-21 -6 (282-4-305-8°K)40 0-824-0-590 (275-5302-l°K)40

d=

3-09(212°K)8

- 2 1 9 rel CF3CO2H (ext)62 - 349 rel CF3CO2H (ext)62 11062

—

—

79ßr, -280-9 siBr, -233-3

BrF 5

1-009108-1-007135 (392-7-446-0°K)4i

e = 4109-0198/ (/ = 0-42°C)4i

30-8-28-2 (291 -6-311 -4°K)40 2-490-1-686(292-1313-2°K)40

3-961-3-678 (77-15-273-15°K)5i 3-263-3-031 (283-40-343-96°K)5i

-173-6 rel SiF463 -222-2 rel SiF463 85063

+ 3 0061

1271,+1073

1271, +1056-6

IF 5

\

1-75(298°K)56

— —

d= 2-7918-0-0049/ (/ = 7-13-24-32°C)56

3-62(128°K)58

—

-334relSiF 4 63

-4-5661

1271, - 1 4 8

—

IF 7

—

0-37-1-25x10-9 (193-256°K)53 - 4 5 1 (298°K)44 15-41 (1,298°K)8 15-48 (g,298°K)8

7-8-9-91x10-8 (213-298°K)65 -58-l(~298°K)44 19-17 (g,298°K)8

5-4xl0-6(298°K)8

—

17-46 (1,298°K)56

<10-9(~298°K)56

1

♦Calculated value. R, Raman-active. i.r., infrared-active. J. W. Mellor, A Comprehensive Treatise on Inorganic and Theoretical Chemistry, Vol. II, Longmans, Green and Co., London (1922). 2 Supplement to Mellor's Comprehensive Treatise on Inorganic and Theoretical Chemistry, Supplement II, Part I, Longmans, Green and Co., London (1956). 3 Gmelins Handbuch der Anorganischen Chemie, 8 Auflage: "Brom", System-nummer 7, Verlag Chemie (1931); "Iod", System-nummer 8, Verlag Chemie (1933). 4 Gmelins Handbuch der Anorganischen Chemie, 8 Auflage: "Chlor", System-nummer 6, Teil B, Lieferung 2, p. 543. Verlag Chemie (1969). 5 N. N . Greenwood, Rev. Pure Appl. Chem. 1 (1951) 84. 6 E. H. Wiebenga, E. E. Havinga and K. H. Boswijk, Adv. Inorg. Chem. Radiochem. 3 (1961) 148. 7 A. G. Sharpe, Non-aqueous Solvent Systems (ed. T. C. Waddington), p. 285. Academic Press (1965). 8 L. Stein, Halogen Chemistry (ed. V. Gutmann), Vol. 1, p. 133. Academic Press (1967). 9 H. Meinert, Z. Chem. 7 (1967) 41. io Table of Atomic Weights, 1969, IUPAC Commission on Atomic Weights: Pure Appl. Chem. 21 (1970) 95. ii Tables of Constants and Numerical Data, No. 17, Spectroscopic Data relative to Diatomic Molecules (general ed. B. Rosen), Pergamon (1970). 12 (a) A. L. McCleMan, Tables of Experimental Dipole Moments, Freeman, San Francisco (1963); (b) E. Herbst and W. Steinmetz, / . Chem. Phys. 56 (1972) 5342. 13 A. G. Gaydon, Dissociation Energies, 3rd edn., Chapman and Hall, London (1968). 14 National Bureau of Standards Technical Note 270-3, U.S. Government Printing Office, Washington (1968). 15 G. F. Calder and W. F . Giauque, / . Phys. Chem. 69 (1965) 2443. 16 R. H. Lamoreaux and W. F . Giauque, / . Phys. Chem. 73 (1969) 755. 17 JANAF Thermochemical Tables, The Dow Chemical Company, Midland, Michigan (1960-8). 18 Adiabatic ionization potentials: S. Evans and A. F . Orchard, Inorg. Chim. Acta, 5 (1971) 81; A. W. Potts and W. C. Price, Trans. Faraday Soc. 67 (1971) 1242. 19 A. P. Irsa and L. Friedman, / . Inorg. Nuclear Chem. 6 (1958) 77. 20 K. H. Boswijk, J. van der Heide, A. Vos and E. H. Wiebenga, Acta Cryst. 9 (1956) 274. 21 G. B. Carpenter and S. M. Richards, Acta Cryst. 15 (1962) 360. 22 L. N . Swink and G. B. Carpenter, Acta Cryst. B24 (1968) 429. 23 H. Stammreich, R. Forneris and Y. Tavares, Spectrochim. Acta, 17 (1961) 1173. 24 E. A. C. Lücken, Nuclear Quadrupole Coupling Constants, p. 288, Academic Press (1969); E. Tiemann, J. Hoeft and T. Törring, Z. Naturforsch. 28a (1973) 1405. 25 R. S. Yamasaki and C. D . Cornwell, / . Chem. Phys. 30 (1959) 1265.

Magnetic susceptibility of the liquid, X ( X106 c g s units) Molar refraction (cm3 mol" 1 )

Specific conductivity of the liquid (ohm - 1 c m - 1 )

cont.)

26 M . P a s t e r n a k a n d T . S o n n i n o , / . Chem. Phys. 48 (1968) 1997. 27 F . W . G r a y a n d J . D a k e r s , Phil. Mag. [7] 11 (1931) 8 1 . 28 H . Selig, H . H . Claassen a n d J. H . H o l l o w a y , / . Chem. Phys. 52 (1970) 3517. 29 R . A . F r e y , R . L . R e d i n g t o n a n d A . L . K . Aljibury, / . Chem. Phys. 54 (1971) 3 4 4 . 30 R . C . K i n g a n d G . T . A r m s t r o n g , / . Res. Nat. Bur. Stand. 74A (1970) 769. 31 M . Schmeisser, W . Ludovici, D . N a u m a n n , P . S a r t o r i a n d E . Scharf, Chem. Ber. 101 (1968) 4 2 1 4 . 32 R . D . B u r b a n k a n d F . N . Bensey, / . Chem. Phys. 2 1 (1953) 6 0 2 . 33 R . D . B u r b a n k a n d F . N . Bensey, j u n . , / . Chem. Phys. 2 7 (1957) 982. 34 K. H . Boswyk and E . H . W i e b e n g a , ^ c t a Cryst. 7 (1954) 4 1 7 . 35 (a) M . S c h m e i s s e r , D . N a u m a n n and E. L e h m a n n , / . Fluorine Chem. 3 ( 1 9 7 3 / 7 4 7 ) 4 4 1 ; (b) R . F o r n e r i s , J. H i r a i s h i , F . A . M i l l e r a n d M . U e h a r a , Spectrochim. Acta, 26A (1970) 581. 36 J . C . E v a n s a n d G . Y . - S . L o , Inorg. Chem. 6 (1967) 836. 37 L . G . A l e x a k o s a n d C . D . Cornwell, / . Chem. Phys. 41 (1964) 2 0 9 8 . 38 J . W . Emsley, J. Feeney a n d L . H . Sutcliffe, High Resolution Nuclear Magnetic Resonance Spectroscopy, Vol. 2 , p . 8 7 1 . P e r g a m o n (1966). 39 A . A . B a n k s , A . D a v i e s a n d A . J . R u d g e , / . Chem. Soc. (1953) 7 3 2 . 40 M . T . R o g e r s a n d E . E . G a r v e r , / . Phys. Chem. 6 2 (1958) 952. 4i M . T . R o g e r s , H . B . T h o m p s o n a n d J . L . Speirs, / . Amer. Chem. Soc. 76 (1954) 4 8 4 1 . 42 M . T . R o g e r s , R . D . Pruett a n d J . L . Speirs, / . Amer. Chem. Soc. 77 (1955) 5280. 43 H . H . H y m a n , T . Surles, L . A . Q u a r t e r m a n a n d A . I . P o p o v , / . Phys. Chem. 7 4 (1970) 2 0 3 8 . 44 M . T . R o g e r s , M . B . Panish a n d J . L . Speirs, / . Amer. Chem. Soc. 7 7 (1955) 5292. 45 A . G . R o b i e t t e , R . H . Bradley a n d P . N . Brier, Chem. Comm. (1971) 1567. 46 W . J . A d a m s , H . B . T h o m p s o n a n d L . S. Bartell, / . Chem. Phys. 5 3 (1970) 4 0 4 0 . 47 E . W . Kaiser, J . S. M u e n t e r , W . K l e m p e r e r a n d W . E . F a l c o n e r , / . Chem. Phys. 5 3 (1970) 5 3 . 48 G . M . Begun, W . H . Fletcher a n d D . F . S m i t h , / . Chem. Phys. 4 2 (1965) 2236. 49 K . O . Christe, E . C . Curtis, C . J. Schack a n d D . Pilipovich, Inorg. Chem. 11 (1972) 1679. 50 H . Selig a n d H . H o l z m a n , Israeli. Chem. 7 (1969) 4 1 7 ; L . E . A l e x a n d e r a n d I . R . Beattie, / . Chem. Soc. (A) (1971) 3 0 9 1 . 5 1 D . W . O s b o r n e , F . Schreiner and H . Selig, / . Chem. Phys. 54 (1971) 3790. 52 H . H . C l a a s s e n , E . L . G a s n e r a n d H . Selig, / . Chem. Phys. 4 9 (1968) 1 8 0 3 ; H . H . Eysel a n d K . Seppelt, ibid. 5 6 (1972) 5 0 8 1 . 53 H . H . R o g e r s , M . T . C o n s t a n t i n e , J . Q u a g l i n o , j u n . , H . E . D u b b a n d N . N . O g i m a c h i , / . Chem. Eng. Data, 13 (1968) 3 0 7 ; W . R . Bisbee, J. V. Hamilton, J. M . G e r h a u s e r a n d R. R u s h w o r t h , ibid. p . 382. 54 R . B o u g o n , J . Chatelet a n d P . Plurien, Compt. rend. 2 6 4 C (1967) 1747. 55 D . Pilipovich , W . M a y a , E . A . L a w t o n , H . F . B a u e r , D . F . S h e e h a n , N . N . O g i m a c h i , R . D . Wilson, F . C . G u n d e r l o y , j u n . , a n d V. E . Bedwell, Inorg Chem. 6 (1967) 1918. 56 H . Selig, C . W . Williams a n d G . J. M o o d y , / . Phys. Chem. 71 (1967) 2739. 57 C . J. Schack, D . Pilipovich, S. N . C o h z a n d D . F . S h e e h a n , / . Phys. Chem. 7 2 (1968) 4697. 58 R . D . B u r b a n k a n d F . N . Bensey, j u n . , / . Chem. Phys. 2 7 (1957) 9 8 1 ; R . D . B u r b a n k , Acta Cryst. 15 (1962) 1207; J. D o n o h u e , Acta Cryst. 18 (1965) 1018. 59 R. D . Burbank and G. R. Jones, Inorg. Chem. 13 (1974) 1071; G. R. Jones, R. D . B u r b a n k and N . Bartlett, ibid. 9 (1970) 2264. 60 M . J . Whittle, R . H . Bradley a n d P . N . Brier, Trans. Faraday Soc. 67 (1971) 2 5 0 5 ; R . H . Bradley, P . N . Brier a n d M . J . W h i t t l e , Chem. Phys. Letters, 11 (1971) 192. 6i S. B u k s h p a n , C . Goldstein, J. S o r i a n o a n d J. S h a m i r , / . Chem. Phys. 5 1 (1969) 3976. 62 E . L . Muetterties a n d W . D . Phillips, / . Amer. Chem. Soc. 81 (1959) 1084; Adv. Inorg. Chem. Radiochem. 4 (1962) 2 4 5 . 63 N . Bartlett, S. B e a t o n , L . W . Reeves a n d E . J . Wells, Canad. J. Chem. 4 2 (1964) 2 5 3 1 . 64 M . T . R o g e r s , R . D . P r u e t t , H . B . T h o m p s o n a n d J . L . Speirs, / . Amer. Chem. Soc. 7 8 (1956) 4 4 . 65 M . T . R o g e r s , J . L . Speirs a n d M . B . P a n i s h , / . Amer. Chem. Soc. 7 8 (1956) 3288.

TABLE 93 {Footnotes

INTERHALOGEN COMPOUNDS

1501

The details of Table 93 disclose that iodine monochloride, monobromide and trichloride are solids at room temperature, whereas bromine trifluoride and pentafluoride and iodine pentafluoride are liquids, and the fluorides of chlorine and iodine heptafluoride are gases. The following sequences of volatility are apparent: C1F > C1F3 < CIF5 BrF 3 < BrF5 IF 5 < IF 7 CIF5 > BrF5 > IF 5

It is noteworthy that the trifluorides CIF3 and BrF3 are significantly less volatile than the corresponding pentafluorides, while IF 7 represents the most volatile interhalogen derivative of iodine. In the solid or liquid states, the materials exhibit electrical conductivities in the range 10~2 to 10 _ 9 ohm - 1 cm - 1 ; the significance of the comparatively high conductivities of some of the liquids is discussed subsequently in connection with the function of these interhalogens as solvent systems. In all cases, however, the electrical conductivity falls below the range characteristic of fused salts. Thus, all of the physical properties of the interhalogens are compatible with their formulation as molecular compounds. Accordingly, variations in the boiling and melting points and in the volatility of the compounds follow the pattern set by molecular polarizability (which is also a function of molecular geometry) and dipole moment (which spans the range from virtually zero for IF 7 to 2-18 D for IF5). The properties of the interhalogens signify that the mononuclear unit XY» is commonly the only recognizable molecular aggregate in the solid, liquid or gaseous phases. A notable exception is provided, however, by iodine trichloride, crystals of which are composed of planar I2CI6 molecules (see below); the mononuclear unit ICI3 has not been characterized. Again, polymeric networks, which denote appreciable interaction between the IX units, have been established for crystalline iodine(I) chloride and bromide. According to the Trouton constants given in the table, chlorine(I) fluoride and iodine(I) chloride would appear also to be associated in the liquid state, but the experimental results may well be substantially in error. The non-ideality of the vapours of chlorine and bromine trifluorides have been attributed to the equilibrium 2XF3^-X2F6

for which equilibrium constants have been reported838; thus, for the chlorine compound, Kp(= p2ciFjPcitFt) is 35-4 atm at 24-2°C. When chlorine or bromine trifluoride is isolated at comparatively high concentrations in solid inert matrices maintained at 5-25°K, the infrared spectrum bears direct witness that dimers and other polymeric species exist as distinct, though weakly associated, aggregates875. Some degree of aggregation is also indicated by the comparatively large Trouton constants reported for the trifluorides. Likewise, a comparison of the Raman spectra of gaseous and liquid iodine pentafluoride points to a measure of association in the liquid876. Contrary to earlier reports, however, the Trouton constant of iodine pentafluoride has now been established as well within the range observed for normal liquids870; evidently, therefore, the degree of association at the boiling point is slight, though this does not invalidate the evidence for association at lower temperatures. 875 R . A . Frey, R. L. Redington and A. L. K. Aljibury, / . Chem. Phys. 54 (1971) 344. 876 H . Selig and H. Holzman, Israeli. Chem. 7 (1969) 417.

1502

CHLORINE, BROMINE, IODINE AND ASTATINE: A. J. DOWNS AND C. J. ADAMS

The only pseudohalide derivatives for which detailed physical constants, thermodynamic properties and structural data have been described are the cyanides C1CN, BrCN and ICN832,837,862,877-879. According to the vibrational-rotational and microwave spectra of the gaseous materials, the molecules are invariably linear with the following dimensions: C1CN, r(C-Cl) = 1-629Ä, r(C = N) = 1-163 Ä; BrCN, r(C-Br) = 1-790Ä, r(C^N) = 1-158 A; ICN, r(C-I) = 1-995Ä, r(C=N) = 1-158 A. The molecular crystals are composed of linear chains · · · X-C = N · · · X - C ^ N · · · , in which the halogen and nitro­ gen atoms of each XCN molecule engage in relatively short-range intermolecular interactions^. The vibrational spectra of the species XCN880, XSCN880 and XN388I have also been reported, and those of the cyanides have been exploited in investigations of the coordinating properties of the molecules, e.g. with respect to various inorganic halides880,882.

2. Thermodynamic properties. Selected thermodynamic functions, based on the most recent estimates838»878'883, appear for the majority of the interhalogen compounds listed in Table 93; calorimetric data are available only for IC1884.885, I2C16885, C1F3827, BrF3»27 and IF 5 870 . The results of the table show that the standard entropies of all the gaseous diatomic molecules form a well-graded series, with the values for the interhalogen compounds falling between those of the parent halogens. The entropy of the unsymmetrical molecule XY exceeds the mean of those for the corresponding symmetrical molecules by a small margin represented, in the ideal case, by R loge2 = 1 - 4 eu. It follows that the entropy change accompanying the gaseous reaction X2+Y2->2XY is small and that the stability of XY with respect to its elements is effectively determined by the appropriate enthalpy change. However, a reaction of the type X2 + HY2 -> 2XY n (n = 3, 5 or 7)

is possible only in the event of a significant negative enthalpy change, since the decrease in the number of molecules occasions a marked decrease in entropy. The most precise estimates of the heats of formation of the diatomic molecules C1F, BrF, IF, IC1 and IBr derive from the convergence limits of the band systems observed in the visible region of the spectrum and thought to arise from the transition ν4 3 Π 0 + <- ΧιΣ + (C1F, BrF or IF) or Amx <- Χ^Σ + (IC1 or IBr), X referring to the ground state of the molecule827»831 »838. The convergence limits thus found are (in cm - 1 ): C1F, 21,512; 877 Tables of Interatomic Distances and Configuration in Molecules and Ions, Chemical Society Special Publication No. 11, London (1958); Supplement, Chemical Society Special Publication No. 18, London (1965). 878 F . D . Rossini, D . D . W a g m a n , W . H . Evans, S. Levine a n d I. Jaffe, Selected Values of Chemical Thermodynamic Properties, Circular 5 0 0 , N a t i o n a l Bureau of S t a n d a r d s , Washington (1952); N a t i o n a l B u r e a u of S t a n d a r d s Technical N o t e 2 7 0 - 3 , U . S . G o v e r n m e n t Printing Office, Washington (1968). 879 A . F . Wells, Structural Inorganic Chemistry, 3rd e d n . , C l a r e n d o n Press, Oxford (1962). 880 H . Siebert, Anwendungen der Schwingungsspektroskopie in der Anorganischen Chemie, SpringerVerlag (1966); K. Nakamoto, Infrared Spectra of Inorganic and Coordination Compounds, 2nd edn., WileyInterscience, New York (1970). 881 D. E. Milligan and M. E. Jacox, / . Chem. Phys, 40 (1964) 2461. 882 K . K a w a i a n d I. K a n e s a k a , Spectrochim. Acta, 25A (1969) 263. 883 w . H . E v a n s , T . R . M u n s o n a n d D . D . W a g m a n , / . Res. Nat. Bur. Stand. 55 (1955) 147; Thermochemical Tables, T h e D o w Chemical C o m p a n y , M i d l a n d , Michigan (1960-8). 884 G . V. Calder a n d W . F . G i a u q u e , / . Phys. Chem. 69 (1965) 2 4 4 3 . 885 R . H . L a m o r e a u x a n d W . F . G i a u q u e , / . Phys. Chem. 7 3 (1969) 755.

JANAF

INTERHALOGEN COMPOUNDS

1503

BrF, 21,190; IF, 23,570; ICl, 17,357; IBr, 14,660. The behaviour resembles closely that of the parent halogen molecules, except that the correlations are rather different as the g or u symmetries of the states do not have to be taken into account. The assumption that the products of dissociation of thefluoridemolecules are X(2P3/2) + F(2^i/2) (X = Cl, Br or I), whereas those of ICl and IBr are normal halogen atoms in their electronic ground states, then leads to the dissociation energies given in Table 93. Some ambiguity still clothes the products of dissociation of the halogen fluorides, for the alternative combination X(2P1/2) + F(2P3/2) cannot be excluded; the corresponding dissociation energies are then consistently lower than those given in the table, viz. C1F, 59-0; BrF, 50-05; IF, 45-6 kcal mol-i. However, from a comparison of bond energies, Slutsky and Bauer886 conclude that the higher dissociation energy is the logical choice for each of the diatomic halogen fluorides. Where thermochemical data are available, as for C1F, BrF, ICl and IBr, the results provide satisfactory support for the suggested modes of dissociation and for the spectroscopic values of the dissociation energy. Combination of the heat of formation of C1F, determined by direct calorimetric measurements, with the spectroscopic value for Z>o°(ClF) implies that £>o0(F2) = ca. 36-5 kcal mol -1 , a conclusion of some importance in that it provided the first cogent evidence against the values of 60-70 kcal at one time favoured for this quantity830»831. By contrast, no corresponding band system has been observed for BrCl; in this case, the thermodynamic parameters recommended by the National Bureau of Standards878 are founded on the experimentally measured equilibrium constants and enthalpy change characterizing the process Br 2 (g)+Cl 2 (g)^2BrCl(g)

these data being supplemented by statistically calculated entropy values. The enthalpies of the reactions and

ClF 3 (g)+2H 2 (g)+ 100H 2 O(l) -► [HCl,3HF,100H 2 O](l) iCl 2 (g)+iH 2 (g)+[3HF,100H 2 O](l) -> [HCl,3HF,100H 2 O](l)

have recently been measured directly in a flame calorimeter at 1 atm and 303-5°K887; together with data from previous investigations, the values yield Ai70/,298[ClF3(g)] = — 39*35 ± 1-23 kcal mol _1 . As applied to the reactions of gaseous CIF5 with hydrogen and ammonia888, the methods of reaction calorimetry furnish A//°/t298[ClF5(g)] = —60-9 ±4-5 kcal mol -1 , while measurements of the equilibrium constant for the reaction ClF 3 (g)+F 2 (g)^ClF 5 (g)

at elevated temperatures suggest a value of -57-7 ± 1 kcal mol- 1 889. For BrF3 andBrF5, the enthalpies of formation have been determined by direct combination of the elements in an adiabatic calorimeter, with the results given in Table 93. The corresponding term favoured for IF 5 is that measured by Settle, O'Hare, Hubbard and Jeffes and referred to in the context of recent calorimetric studies of the pentafluoride870. From equilibriumpressure measurements of the dissociation IF 7 (g)^IF 5 (g)+F 2 (g) 886 L. Slutsky and S. H. Bauer, / . Amer. Chem. Soc. 76 (1954) 270. 887 R. c . King and G. T. Armstrong, / . Res. Nat. Bur. Stand. 74A (1970) 769. 888 w . R. Bisbee, J. V. Hamilton, J. M. Gerhauser and R. Rushworth, / . Chem. and Eng. Data, 13 (1968) 382. 889 R. Bougon, J. Chatelet and P. Plurien, Compt. rend. 264C (1967) 1747.

1504

CHLORINE, BROMINE, IODINE AND ASTATINE: A. J. DOWNS AND C. J. ADAMS

Bernstein and Katz calculate that ΔΗ° = 28-5 ±2-0 kcal mol"* and AS° = 43-5 ± 3-0 eu in the temperature range 45O-550°C838. On the basis of their measurements and of the enthalpy of formation of IF5, ΔΗ°/,29ζ[1¥Ί(ζ)] is -229-9 kcal mol-i. Mean bond energies for the interhalogen molecules appear in Table 93, while, for the halogenfluoridesXFn, the variation of bond energy as a function of X and n is depicted in Fig. 34. For a given molecular type XFn, the general sequence of bond energies is /U 1

i

1

1 IF

1

I

I

I

Br

I

60

50

40

in

J Ci

FIG. 34. Average bond energies of halogen fluorides.

ClFn < BrFw < IFn, the energy increasing as the X-F bond becomes more polar; CIF and BrF form the only exceptions to this rule. The bond energies of the diatomic interhalogen molecules are reproduced satisfactorily by the Pauling relationship between electronegativity X and bond energy BS90: (XX-XY) 2 = ~ [ l ? ( X - Y ) -

VBQC-X)B(Y-Y))

For a given halogen X, the X-F bond energy of XFn diminishes as n increases, as in the series CIF, C1F3, C1F5, but the rate of decrease of bond energy with respect to n is notably less when X = I than when X = Cl. To reproduce the bond energies of polyatomic interhalogen molecules XFn by means of the Pauling relationship, it is necessary to assign 890

D. A. Johnson, Some Thermodynamic Aspects of Inorganic Chemistry, p. 153, Cambridge (1968).

1505

INTERHALOGEN COMPOUNDS

to X an effective electronegativity which increases in step with w. From the bond energies of the other halogen fluorides (Fig. 34), the average bond energy of IF 3 is estimated to be ca. 66 kcal; the enthalpy of formation of the gaseous compound from the elements in their standard states at 25°C is then ca. — 116 kcal mol _1 . The stability of the gaseous diatomic molecules XY with respect to the gaseous elements is determined almost exclusively by the difference between the bond energy of the XY molecule and the mean of those of the X 2 and Y 2 molecules; in unison with this energy term, which reflects the difference in electronegativity between X and Y, the stability decreases in the order IF > BrF > C1F > IC1 > IBr > BrCl. The gain in energy with the formation of an interhalogen XY may thus be ascribed to the polar character of the X-Y bond. With respect to dissociation into the gaseous elements, all of the interhalogen molecules are thermodynamically stable at 25°C, though only marginally so in the case of BrCl. However, the existence of a particular interhalogen compound XYW, even when it contains a strong X-Y bond, depends also on its thermodynamic stability with respect to other compounds containing the atoms X and Y, and on the speed of whatever disproportionation processes may be open to it. A plot of AG°/,298 f° r gaseous halogen fluoride molecules XFn as a function of n (Fig. 35) gives clear notice, through the relative slopes AG°/An, that, whereas C1F is stable, BrF and IF are thermodynamically vulnerable to disproportionation +1 1

Oxidation state of X +5 +3 1

1

+7 1

1

^^C1F5

-50

\5

rF

3

^NgrFj -100

V

a

1 o

< -150

\i -200 h

^*·^

-250 FIG. 35. Free energies of formation (at 298°K) of gaseous halogen fluorides XF n .

1506

CHLORINE, BROMINE, IODINE AND ASTATINE! A. J. DOWNS AND C. J. ADAMS

processes of the type 3XF->

X2+XF3

and 5XF->2X2+XF5

None of the other diatomic interhalogen molecules, BrCl, ICl or IBr, shows any tendency to undergo reactions analogous to these. The qualitative evidence presently available854 suggests that, in the condensed phase, IF3 disproportionates at low temperatures in accord­ ance with the reaction 2IF3->IF+IF5

or, at room temperature, in accordance with the reaction 5IF3-*I2+3IF5

In general, disproportionation of intermediate interhalogen compounds XYn is likely to occur in those situations where the bond energy decreases but little with increase of n (cf. IF, IF 3 and IF 5 in Fig. 34). A notable feature of the thermodynamic properties of the interhalogen compounds is the unique stability of iodine pentafluoride (see Table 93 and Fig. 35), which is not disposed to disproportionate and may be expected to dissociate significantly only at temperatures in excess of 1000°C, e.g. ~1400°C

IF 5 (g)

► IF 3 (g)+2F(g)838

In this respect there is evidently a close parallel in thermodynamic status between IF 5 and the oxy-species I2O5 and IO3-. At elevated temperatures, the marked increase of entropy accompanying reactions such as 350° O

ClF 5 (g) ,

ClF 3 (g)+F 2 (g)889

>550°c BrF 5 (g)

and

► BrF 3 (g)+F 2 (g)838

>45o°c lF7(g)

► IF 5 (g)+F 2 (g)838

promotes the dissociation of certain polyatomic halogen fluorides, while even at relatively low temperatures the reaction ICl3(g)->ICl(g)+Cl2(g) appears to proceed effectively to completion. 3. Structural properties829,835,838-840,877,879,89i,892. As is evident from the details of Table 93 and Fig. 36, the structures of almost all the interhalogen compounds are now known, the only exceptions being C1F5, for which no definitive structural parameters are yet available, and the ill-defined IF3. The most precise estimates of the interatomic distances in the diatomic molecules come from the microwave spectra and from the rotational detail observed at high resolution in the electronic spectra of the gaseous materials891. For 891 Tables of Constants and Numerical Data, N o . 17, Spectroscopic (general e d . B . Rosen), Pergamon (1970). 892 H . A . Bent, Chem. Rev. 68 (1968) 599.

Data relative to Diatomic

Molecules

INTERHALOGEN COMPOUNDS 893

1507

894

C1F3 and BrF 3 microwave studies of the vapours again provide definitive structural details, while the combination of microwave and electron-diffraction measurements has recently enabled the structures of BrF 5 and IF5 to be determined with some confidence895. Electron diffraction also forms the basis of the most detailed structural analysis so far carried out on the IF 7 molecule in the vapour phase896. The geometry of C1F5 has been deduced from its vibrational and 1 9 Fnmr spectra; though consistent with the vibrational properties, the approximate dimensions given in the table are prompted largely by analogy with related molecules. X-ray diffraction measurements have been reported for numerous interhalogen com­ pounds in the crystalline condition, but only in the cases of IC1897, IBr898 and I 2 C1 6 899 have the results been refined to obtain closely defined structural parameters; less well characterized, with inferior structure factors, are crystalline CIF3900, BrF3901, BrF 5 901 and IF790i,902. Crystalline IF 5 is monoclinic with a structure so complicated as to militate against significant refinement of the raw crystallographic data; however, the recently determined crystal structure of the molecular addition compound XeF 2 ,IF 5 discloses the presence of more-or-less discrete IF 5 molecules, whose structure has thus been established with some precision903. The measured interatomic distance in a diatomic interhalogen molecule XY is generally somewhat shorter than the sum of the covalent radii of X and Y (Fig. 36), the margin increasing as the electronegativities of X and Y diverge. Such a pattern complies at once with the variation in bond energies for this group of molecules, and with the general implica­ tions of the empirical Schomaker-Stevenson equation relating bond length to electronega­ tivity difference879. It is now established beyond dispute that the C1F3 and BrF 3 molecules have planar skeletons of C2v symmetry with an unusual T-shaped configuration. Each molecule contains two nearly collinear X-F bonds, which are, to a first approximation, perpendicular to the third, though the idealized T-form is distorted in such a way that the angle between the bonds is slightly, but significantly, less than 90°. The unique X-F bond is somewhat shorter, whereas the other two bonds are somewhat longer, than that in the corresponding diatomic molecule. The fluorine atoms of the BrF 5 and IF 5 molecules assume the form of a square-based pyramid belonging to the symmetry group C&, the bromine or iodine atom being situated somewhat below the basal plane of the pyramid, so that the angle between the apical and basal X-F bonds is rather less than 90°. In each of these molecules, the unique, apical bond is shorter than those of the basal XF4 unit, but, whereas the limits defined by the two Br-F distances in BrF 5 span that in diatomic BrF, all the bonds in IF 5 are appreciably shorter than that in IF. Curiously, too, with respect to idealized octahedral coordination with angles of 90° at the central atom, IF 5 experiences the greater angular distortion (^F a p l c a l -I-F b a s a l = 81-9°; ZF a p i c a l -Br-F b a s a l = 84-8°), but with a 893 D . F . Smith, / . Chem. Phys. 21 (1953) 609. 894 D . W. Magnuson, / . Chem. Phys. 27 (1957) 223. 895 A. G. Robiette, R. H. Bradley and P. N . Brier, Chem. Comm. (1971) 1567. 896 w . J. Adams, H. B. Thompson and L. S. Bartell, / . Chem. Phys. 53 (1970) 4040. 897 K. H. Boswijk, J. van der Heide, A. Vos and E. H. Wiebenga, Acta Cryst. 9 (1956) 274; G. B. Carpenter and S. M. Richards, ibid. 15 (1962) 360. 898 L. N . Swink and G. B. Carpenter, Acta Cryst. B24 (1968) 429. 899 K. H. Boswijk and E. H. Wiebenga, Acta Cryst. 7 (1954) 417. 900 R . D . Burbank and F. N . Bensey, / . Chem. Phys. 21 (1953) 602. 901 R. D. Burbank and F . N . Bensey, jun., / . Chem. Phys. 27 (1957) 982. 902 j . Donohue, / . Chem. Phys. 30 (1959) 1618; Acta Cryst. 18 (1965) 1018. 903 G. R. Jones, R. D . Burbank and N . Bartlett, Inorg. Chem. 9 (1970) 2264.

1508

CHLORINE, BROMINE, IODINE AND ASTATINE: A. J. DOWNS AND C. J. ADAMS

(a) Type XY

lolecule CIF BrF IF BrCl IC1 IBr

r

1.62811 1.756 1.9089 2.138 2.3207 2.485

calc.» Ä-

1.703 1.851 2.042 2.136 2.327 2.475

f

e

r

calc.

-0.075 -0.095 -0.133 + 0.002 -0.006 + 0.010

(b) Type XY

Molecule

R, A

C1F3 BrF,

1.698 1.810

1.721

87° 29' 86° 12.6'

(c) Type X2Y6

h<** r.A R, Ä «, °

ß.9

2.38, 2.39 2.68, 2.72 84 94

FIG. 36. Structures of interhalogen molecules of the types XY, XY3 and Χ2Υβ·

1509

INTERHALOGEN COMPOUNDS

(b) jS-ICI

(a) α-ICl

Q

Φι°

1%

?ΐ

I

vb

F

Q

■ιό

ο^κ, φ-ο-ο

JO, * J3-ICI rr Ä r2, Ä R„ Ä

4Ä

2.44 2.37 3.00 3.08

2.440 2.351 2.939 3.060

cf. r e (I-a) = 2-321 Ä; sum of van der Waals radii: r(I) + r(Cl) « 3·95 Ä; 2 x r ( I ) -4-30 Ä. F I G . 36. Structures of a- and j8-forms of crystalline ICl.

1510

CHLORINE, BROMINE, IODINE AND ASTATINE: A. J. DOWNS AND C. J. ADAMS

(a) IBr Arrangement within each mirror plane of the crystal

r, Ä 2.524 R r Ä 3.181 R„, Ä 3.764

α, ß, γ, δ,

° ° ° °

175.2 94.6 104.4 165.5

cf. r e (I-Br) = 2-485 Ä; sum of van der Waals radii: r(I) + r(Br) = 4-10 Ä; 2 x r(I) = 4-30 Ä.

(b) Type XY5

(c) Type XY7

x

\

Molecule BrF5 IF,

r, Ä R, Ä 1.689 1.774 1.844 1.869

IF

'

r, Ä 1.786 R, Ä 1.858 Angles correspond to those of a regular pentagonal bipyramid (symmetry group D5h)

α, ° 84.8 81.9

FIG. 36. Structures (a) of solid IBr, (b) of molecular interhalogens of the type XY5, and (c) of molecular IF7.

INTERHALOGEN COMPOUNDS 895

1511

smaller disparity between apical and basal bond lengths . The isolated IF 7 molecule has the form of a pentagonal bipyramid belonging to the symmetry group D5h, in which the bonds of the apical F-I-F unit are shorter than those of the more congested equatorial IF 5 unit; the mean I-F distance is shorter than that observed in IF5, despite the increased crowding. However, the details of the electron-diffraction pattern896 imply that the molecule is deformed from D5h symmetry on the average by puckering of the equatorial ring of fluorine atoms and by bending of the axial F-I-F unit, the angular displacements about the central atom being 7-5° and 4-5°, respectively. The simplest interpretation of the diffrac­ tion intensities is that the molecule undergoes essentially free pseudorotation, involving facile intramolecular rearrangement, along a pathway (predominantly described by e2" displacement coordinates) connecting 10 equivalent structures of C 2 symmetry via inter­ mediates of Cs symmetry. The implication of deformed, polar configurations in this mechanism accounts for the fact that at low temperatures, but not at room temperature, a molecular beam of IF 7 is deflected by an inhomogeneous electric field, the molecules behaving as if they possess electric dipole moments904. The extraordinary flexibility characteristic of the IF 7 molecule appears also to be reproduced, with heightened ampli­ tudes of distortion, in ReF7, the only other molecule of the type XY 7 which has been analysed in this way905. On the basis of available crystallographic data, the molecules C1F3, BrF3, BrF5 and IF5 suffer comparatively minor changes in the transition from the vapour to the solid phase; likewise, the dimensions reported for the IF 5 molecule in crystalline XeF2,IF5 are in close agreement with those of the gaseous molecule. There has been considerable controversy over the interpretation of X-ray diffraction measurements for IF7, the issue being complicated by the apparent disorder in the crystal. Inevitably, it can now be appreciated, the ease with which the free molecule can be deformed from one configuration to another causes complica­ tions when the molecules pack together in the crystal. The most recent analyses conclude that the X-ray diffraction data are consistent with, but do not prove, the presence of IF 7 molecules with D5h symmetry; a distance of 1-80 A has thus been deduced for each of the seven I-F bonds902. In both IC1897 and IBr898 in the crystalline state, the IX molecules are incorporated in polymeric networks (Fig. 36); that there is relatively strong interaction be­ tween the molecules is evident from the observations (i) that the intermolecular distances are significantly shorter than the sum of the van der Waals' radii of the atoms concerned, and (ii) that the intramolecular I-X distances are somewhat longer than those in the gaseous IX molecules. IC1 crystallizes in two forms, both of which are characterized by zigzag chains containing two types of IC1 molecule: in molecules of one type, both atoms partici­ pate in short-range intermolecular interactions along the chain; in molecules of the other type, the chlorine atoms branch off the chain and are not involved in such interactions. With respect to one another, the branching chlorine atoms are disposed trans in jS-ICl, but eis in α-ICl. Further, the chains in j8-ICl are essentially planar, while in α-ICl they are puckered. The structures suggest that, in the intermolecular interactions along the chains, the chlorine atoms act chiefly as electron donors, while the iodine atoms act either as electronacceptors or simultaneously as donors and acceptors, after the fashion of the iodine atoms in solid iodine892. Whereas there are thus only half as many strong intermolecular interactions per molecule in crystalline IC1 as in iodine, the molecules of IBr crystals are linked to form planar sheets in a herring-bone pattern analogous to that in iodine. Of the 904 E . W. Kaiser, J. S. Muenter, W. Klemperer and W. E. Falconer, / . Chem. Phys. 53 (1970) 53. 90s E . J. Jacob and L. S. Bartell, / . Chem. Phys. 53 (1970) 2235.

1512

CHLORINE, BROMINE, IODINE AND ASTATINE: A. J. DOWNS AND C. J. ADAMS

interhalogen compounds iodine trichloride appears to be unique in that the crystals are composed of planar, symmetrical I2C16 molecules separated from one another by normal van der Waals' distances. Within each I2CI6 molecule, the terminal I-Cl bonds are of about the same length as in IC1, but the bridging I-Cl bonds are considerably longer. 4. Spectroscopic properties. In contrast with a homonuclear halogen molecule, a di­ atomic interhalogen molecule XY lacks a centre of symmetry. Thus, although the valenceshell structure is similar to that of a symmetrical halogen molecule, orbitals may no longer be labelled with a parity suffix. Corresponding atomic orbitals still mix with one another to produce σ- or ττ-molecular orbitals, though the coefficients designating the contributions made by the individual atoms, viz. a and b in αψχ ± bi/jY, are no longer equal. Hence, while the valence shell of an XY molecule may be written as (1σΥ)2(2σΧ)2(3σΥ)2(ΐ7ΓΥ)4(2πΧ)4

1Σ +

where X is the halogen of higher atomic number, mixing interaction increases the extent to which the 2^orbital should be associated with the Y atom. The 1σ and 2σ molecular orbitals derive mainly from the valence j-orbitals of the X and Y atoms, while the remaining high-energy molecular orbitals are functions primarily of the corresponding /?-orbitals. Promotion of a (ITTY)- or (27rX)-electron to the anti-bonding (4aX)-orbital gives rise to the excited Π-states B(l σΥ)2(2σΧ)2(3σΥ)2(17τΥ)3(27τΧ)4(4σΧ)ΐ

I >m

Α(1σΥ)2(2σΧ)2(3σΥ)2(1πΥ)4(2πΧ)3(4σΧ)ΐ

1.3Π

Transitions implicating these and other excited states have been observed, either in absorp­ tion or in emission, for all of the diatomic interhalogen molecules, and physical constants, such as transition energies, vibrational frequencies and internuclear distances, have been deduced for some of the states thus identified. For fuller details the reader is referred elsewhere827,89i,906. Although absorption spectra in the visible and ultraviolet regions have also been described for some of the polyatomic interhalogen compounds838, no proper interpretation of the results has yet been feasible. In common with the parent halogens, gaseous BrCl, IC1 and IBr have been shown to exhibitfluorescence,and for all three molecules resonance Raman and resonance fluorescence spectra have lately been reported907. The electronic absorption spectra of the species BrCl, IC1 and IBr dissolved in a variety of solvents have also been investigated, and, as with the parent halogen molecules, evidence has thus been adduced for charge-transfer interactions between the interhalogen and donor molecules. In the acceptor action thus manifest, the uniformity of behaviour of diatomic halogen molecules is such that it would be unduly artificial to separate homonuclear X 2 from heteronuclear XY species; accordingly, this particular aspect of the behaviour of interhalogen compounds has already been treated in Section 2 (see p. 1196). Recent measurements of the photoelectron spectra of the gaseous molecules IC1 and IBr908 have led to the ionization potentials recorded in Table 93. The spectra show obvious analogies to those of the parent halogen molecules, the observed bands being consistent with the anticipated energy sequence (3σΥ)2 < (ΙττΥ)4 < (2?rl)4 (Y = Cl or Br). Upper 9 °6 G. Herzberg, Molecular Spectra and Molecular Structure. I. Spectra of Diatomic Molecules, 2nd edn., p. 501, van Nostrand (1950). 907 w . Holzer, W. F. Murphy and H. J. Bernstein, / . Chern. Phys. 52 (1970) 399.

INTERHALOGEN COMPOUNDS

1513

limits for the ionization potential have also been based on the appearance potentials of the appropriate molecular ions observed in mass-spectrometric studies of these and other interhalogen compounds909. Numerous studies testify to the interest occasioned by the vibrational properties of the interhalogen molecules838»910»911. The earlier measurements of the vibrational spectra of polyatomic systems had as a primary objective the elucidation of the molecular geometries. Thus, one of the first positive indications of the C4v symmetry of the C1F5, BrF5 and IF 5 molecules was afforded by their infrared and Raman spectra911. However, with the more definitive conclusions of microwave, electron-diffraction and X-ray investigations, the main ambitions of the most recent experiments have been (i) to determine for an individual molecule reliable assignments of the vibrational frequencies, (ii) to explore the nature of the force field encompassing the atoms of the molecule, and (iii) to monitor chemical changes affecting the compound. Evaluated, for the most part, by reference to vibrational progres­ sions in the electronic spectra, the vibrational frequencies of the isolated diatomic molecules signify that the stretching force constant decreases uniformly in the sequence C1F > BrF > IF > BrCl > IC1 > IBr, which complies with the pattern of increasing internuclear distance, but not with the somewhat irregular pattern of bond energies (see Table 93). The infrared and Raman spectra observed for several of these compounds confirm the findings based on the electronic spectra907»910; further, variations of the frequencies, or of other spectroscopic properties, which accompany changes of phase or solvent or the forma­ tion of distinct molecular adducts, serve as an index to the intermolecular interactions experienced by the molecules in the condensed phases. The reduction both of the vibrational frequency and of the stretching force constant defining the X-Y unit gives clear notice of the anti-bonding character of the acceptor orbitals in molecules like IC1. The recently reported Raman spectra of gaseous C1F3 912, BrF3 912, IF 5 876 and IF7 913 , together with the infrared spectra of CIF3, BrF3 and BrF5 in the matrix-isolated condition875, have brought new facts to light about the vibrational spectra of these molecules, as well as eliminating some of the uncertainties and errors in earlier accounts838. The vibrational assignments thus deduced are presented in Table 93. The force fields for the polyatomic halogen fluorides are as yet poorly determined, though normal coordinate analyses of the molecules XF 3 (X = ClorBr)875, XF 5 (X = Cl, Br or I)911 and IF 7 913 yield valence force constants, which clearly discriminate between the two types of X-F bond found in each of these molecules; in every case, the more attenuated X-F bonds are characterized by smaller force constants. It is noteworthy that the longer bonds in BrF3 have higher force constants than the corresponding bonds in C1F3, whereas corresponding parameters for molecules of the type XF 5 increase in the order C1F5 < BrF5 < IF5. As previously noted, the vibra­ tional spectra of the molecules CIF3875, BrF3875 and IF 5 876 under various conditions allude to significant degrees of polymerization in the condensed phases at low temperatures. However, whereas the dimers X2F6 isolated in solid inert matrices at 5-25°K appear to be of C2h symmetry and to involve unsymmetrical X · · · F-X bridges, the infrared and Raman 908 A . W. Potts and W. C. Price, Trans. Faraday Soc. 67 (1971) 1242; S. Evans and A. F. Orchard, Inorg. Chim. Acta, 5 (1971) 81. 909 A. P. Irsa and L. Friedman, / . Inorg. Nuclear Chem. 6 (1958) 77. 910 H. Stammreich, R. Forneris and Y. Tavares, Spectrochim. Acta, 17 (1961) 1173. 9Π G. M. Begun, W. H. Fletcher and D. F. Smith, / . Chem. Phys. 42 (1965) 2236. 912 H. Selig, H. H. Claassen and J. H. Holloway, J. Chem. Phys. 52 (1970) 3517. 913 H. H. Claassen, E. L. Gasner and H. Selig, / . Chem. Phys. 49 (1968) 1803; R. K. Khanna, / . Mol. Spectroscopy, 8 (1962) 134.

1514

CHLORINE, BROMINE, IODINE AND Α8ΤΑΉΝΕ: A. J. DOWNS AND C. J. ADAMS

spectra of solid iodine(III) chloride substantiate the presence of symmetrical I-Cl-I bridges in the planar I2C16 molecules914. Although geometrically analogous to Au2Cl6, I2CI6 differs from this molecule in its abnormally long I-Cl bridge bonds, which are also un­ commonly weak, as judged by their stretching force constants. The 19F nmr spectra of all the halogen fluorides except BrF, IF and IF 3 have been recorded, in some instances in both the liquid and vapour states838; the relevant chemical shifts and coupling constants are collected in Table 93. Although the chemical shifts are relatively widely scattered, the resonances invariably occur at low fields (on a scale which has the F 2 and HF molecules as, respectively, the low- and high-field limits), implying, as with the noble-gas fluorides915, that the fluorine atoms are relatively poorly shielded. The general trend in a series offluoridesXF» is for the shielding of the fluorine to decline as n increases or as the atomic number of X decreases. Because of the inherent difficulties in direct application of Ramsey's theory of the chemical shift, a full interpretation of the factors contributing to the shielding is not yet accessible, though the results can be rational­ ized in terms of a substantial degree of ionic character in the X-F bonds915; numerical estimates of ionic character and of the charge densities at thefluorineatoms have thus been pronounced. It is noteworthy that the apical fluorine atoms of C1F5, BrF5 and IF 5 are less well shielded than the basal atoms; similarly, the two fluorine atoms of the nearly linear F-Cl-F unit of C1F3 are less well shielded than the third fluorine. This reduction in shielding of certain fluorine atoms is attributable to the virtually exclusive use of atomic /^-functions in the bonding of such atoms. Under suitable conditions, the 19F nmr spectra of the mole­ cules CIF3, CIF5, BrF5 and IF 5 give unequivocal notice of the geometries of the molecules, distinct multiplet resonances due to magnetically non-equivalent fluorines being observed. The temperature-dependence of the liquid-phase spectra indicates that intermolecular fluorine-exchange becomes rapid at ca. 60°C for C1F3 and at ca. 200°C for IF5, the estimated activation energies being 4-8 and 13kcalmol -1 , respectively. Fluorine-exchange in CIF3 is much accelerated by the presence of hydrogen fluoride, even at very low concentrations. Thus, although the exchange may well proceed, as first suggested by Muetterties and Phillips916, via transientfluorine-bridgeddimers of the type identified by matrix-isolation, ionization as well as association may also provide a pathway for exchange. An analogous mechanism can be postulated for thefluorine-exchangein IF5. No sign of rapid fluorineexchange has been found with BrF5 at temperatures up to 180°C, implying that the activation barrier to such exchange is significantly greater in this case, but, to judge by the single 19F resonance observed for the liquid,fluorine-exchangebetween BrF3 molecules is rapid even at room temperature. On the nmr timescale, equivalence is preserved for each of the seven fluorine atoms of IF7, irrespective of the phase or temperature, but in this instance the rapid wframolecular rearrangement inferred from other measurements (q.v.) is presumed to be responsible for the averaging of the magnetic environments experienced by the fluorine atoms. The observed temperature-dependence of the spectrum of liquid IF 7 is then attrib­ uted, not to intermolecular exchange, but to the effects of coupling between the 19F and quadrupolar 127I nuclei838. 914 H . Stammreich a n d Y . K a w a n o , Spectrochim. Acta, 24A (1968) 899; R . Forneris, J. Hiraishi, F . A . Miller a n d M . Uehara, ibid. 26A (1970) 581. 915 H . H . H y m a n , Science, 145 (1964) 773; J. C H i n d m a n a n d A . Svirmickas, Noble-gas Compounds (ed. H . H . Hyman), p . 251, University of Chicago Press, Chicago (1963); C . J. Jameson a n d H . S. Gutowsky, / . Chem. Phys. 40 (1964) 2285. 916 E . L . Muetterties a n d W . D . Phillips, / . Amer. Chem. Soc. 79 (1957) 322.

1515

INTERHALOGEN COMPOUNDS

Quadrupole coupling of the heavier halogen nuclei in interhalogen compounds has also been investigated through the agency of the microwave spectra of the vapours or through measurements of the nqr or 129I Mössbauer spectra of the solids917»918. The quadrupole coupling constants and, where known, the asymmetry parameters are included in Table 93, as are the 129I isomer shifts characterizing the Mössbauer spectra of IC1, IBr, I2CI6, IF 5 and IF 7 918 . The significance of these parameters has been discussed in Section 3.2 (pp. 1271-6 and Fig. 22), and estimates of the ionic character of certain interhalogen bonds, based on the classical Townes-Dailey approximation, appear in Table 27. The general tenor of the data supports the view that bonding in interhalogen molecules arises primarily through the interactions of valence /7-electrons of the halogen atoms, with but a meagre contribution from the atomic ^-functions; to judge, for example, by the 129I isomer shifts, ^-orbital participation appears also to be minimal in all but the IF 7 molecule918. Chemical Properties 8 ^ S40

1. General characteristics. In general, the chemical properties of the interhalogen compounds ΧΥΛ resemble those of the parent halogens. The two principal categories of their reactions are: Donor-acceptor {or Acid-base) Interactions These range from the relatively weak intermolecular interactions involving XYn molecules to interactions which may be strong enough, in effect, to bring about the heterolytic dissociation of an X-Y bond. XYn-f D «, * XYn,D

( D - amine or ether)

'Outer' charge-transfer complex XYnZ~ -

2

XY

A

^

χ γ

n-1

+

+ Α γ

-

Z halogen which may be the same as, or different from,

Pure liquid "

A =

^

i(m

acceptor, e.g. SbF

Thus, with donor molecules such as those of amines or ethers, which resist oxidation, an interhalogen molecule like BrCl, IC1 or IBr experiences weak intermolecular binding. The properties of many of the solutions formed by the interhalogen compound in different solvents and of the low-melting addition compounds afforded by some of these systems are explicable in terms of charge-transfer interactions between the molecular components (see Section 2, p. 1196) 826-829 ' 919 . Presumably to be formulated thus are the crystalline complexes formed by neutral oxygen- or nitrogen-bases, not only with a diatomic interhalogen molecule, e.g. BrCl,py (py = pyridine or substituted pyridine)920 and 917 E. A. C. Lücken, Nuclear Quadrupole Coupling Constants, p. 288, Academic Press (1969). 918 M. Pasternak and T. Sonnino, J. Chem. Phys. 48 (1968) 1997; S. Bukshpan, C. Goldstein, J. Soriano and J. Shamir, ibid. 51 (1969) 3976. 919 L . J. Andrews and R. M. Keefer, Adv. Inorg. Chem. Radiochem. 3 (1961) 9 1 ; R. S. Mulliken and W. B. Person, Molecular Complexes, Wiley, N e w York (1969). 920 T . Surles and A. I. Popov, Inorg. Chem. 8 (1969) 2049.

1516

CHLORINE, BROMINE, IODINE AND ASTATINE: A. J. DOWNS AND C. J. ADAMS

IF,quinoline921, but also with bromine pentafluoride922, iodine tri-839,923 a n c j pentafluoride838, and iodine trichloride827»839, e.g. BrF5,pyridine, IF3,pyrazine, IF5,dioxan, IF5,HCONMe2 and ICl3,pyridine, though it is possible that at least some of these materials are better represented by an ionic constitution, e.g. [(quinoline)2I]+IF2~ or [(pyridine)2l]+ IF6-839. Charge-transfer spectra have recently been reported even for iodine heptafluoride in n-hexane or cyclohexane solutions924. According to the crystallographic evidence con­ cerning certain of the crystalline adducts (Table 16), the X-Y bond of the interhalogen molecule is commonly attenuated by the interaction, but the integrity of the molecule is preserved. Compared with a parent halogen molecule, an interhalogen is, by virtue of its polarity, a more potent acceptor; it is also more susceptible to heterolytic cleavage. Potential cationic products of such a process which have lately been identified with some degree of assurance include XF 2 + and XF 4 + (X = Cl, Br or I), C1 2 F + , IC12+ and IF 6 + (see Section 4A)925, while numerous polyhalide anions, for example of the type XY n Z~ (Z = a halogen which may be the same as, or different from, X or Y), are now well authenticated (see p. 1534)835»840»841. In accordance with the behaviour outlined in the scheme above, most interhalogen compounds are ambivalent with respect to halide ions, which they have the capacity to accept in forming a polyhalide anion or to relinquish in forming a cationic derivative. Heterolytic cleavage of an interhalogen bond is encouraged by a powerful halide ion-acceptor, e.g. BF3, AsF 5 or SbX5 (X = F or Cl), and by conditions tending to stabilize the highly electrophilic interhalogen cation thus formed, which can exist only in the presence of molecules and anions of very low basicity. Conversely, poly­ halide anions are stabilized by environments low in acidity. Halogenation Here heterolytic or homolytic cleavage of an interhalogen bond effects halogenation of a reagent. Thus, diatomic interhalogen molecules add to unsaturated units, e.g. XY

H-

CIF

+

\:=c

_£

£

I

I

X SF4

827,926,927

Y

SF5C1 928

Substitution reactions, often with simultaneous oxidation of the reagent, form a chemical common denominator for interhalogen compounds. Typical of such behaviour is the fluorinating action of halogen fluorides on a wide range of elements and their binary derivatives (see p. 1525)838. 2. Acid-base properties: solvent functions836»838. Because of their great reactivity, the interhalogens dissolve only a limited number of substances without causing them to 921 M . Schmeisser, P . Sartori and D . N a u m a n n , Ber. 103 (1970) 880. 922 H . Meinert and U . Gross, Z. Chem. 9 (1969) 190. 923 M . Schmeisser, K . D a h m e n and P. Sartori, Ber. 103 (1970) 307; M . Schmeisser, P . Sartori a n d D . N a u m a n n , ibid. p p . 312, 590. 924 p . R . H a m m o n d , / . Phys. Chem. 74 (1970) 647. 925 R . j . Gillespie and M . J. Morton, Quart. Rev. Chem. Soc. 25 (1971) 553. 926 w . K . R . Musgrave, Adv. Fluorine Chemistry, 1 (1960) 1. 927 R . D . Chambers, W . K . R . Musgrave and J. Savory, Proc. Chem. Soc. (1961) 113; / . Chem. Soc. (1961) 3779; P . Sartori a n d A . J. Lehnen, Ber. 104 (1971) 2813. 928 F . N y m a n a n d H . L. Roberts, / . Chem. Soc. (1962) 3180

INTERHALOGEN COMPOUNDS

1517

undergo reactions more drastic than ion-formation or solvation. Most investigations of these substances as solvents have been purely qualitative in nature, and, even where quantita­ tive data are available, their interpretation is often obscure. There are many practical obstacles to such investigations: BrF3, the only interhalogen solvent to have the status of a reagent of some importance in preparative inorganic chemistry, explodes with water and most organic matter, reacts with asbestos with incandescence, and can be manipulated in quartz apparatus only because of the sluggishness of the reaction with silica in this form at normal temperatures; the iodine chlorides, though they are less violent in their reactions, nevertheless possess the property, inconvenient in electrochemical studies, of readily dissolving gold and platinum. Measurements of electrical conductivity provide some indica­ tion of the number and mobility of the ions present, but give little information as to their nature. Identification of products liberated at electrodes may throw some light on this problem, but too often the relationship between what carries the current and what is discharged at the electrodes is open to question. Even if, as is sometimes the case, the struc­ tures of solid adducts derived from the solvent system are known, far-reaching changes may occur on dissolution, and halogen-exchange processes are often very rapid. Transport and emf measurements, which have contributed so much to our knowledge of the nature of aqueous solutions, have not been performed to any profitable extent in interhalogen solvents. The picture presented by experimental findings is therefore not a very clear one, and much further work will be needed before the nature of the solutions is well understood. Intimately related to the solvent properties of an interhalogen compound are the reac­ tions whereby a halide ion is formally gained or lost (q.v.), since these represent acid-base functions characteristic of the compound. Of particular relevance therefore is the behaviour (i) with respect to a Lewis acid like BF3 or SbCl5 endowed with the power to abstract a halide ion, and (ii) with respect to a base which can serve, directly or indirectly, as a source of halide ions. Where solid complexes arise from such interactions (Table 94), characteriza­ tion has been achieved, sometimes definitively by X-ray crystallographic investigations, more often in terms of the vibrational or nmr spectra of the materials or of their electrical conductivities in solution. However, the acid or base strength of an interhalogen molecule is only one of several factors that contribute to the effectiveness of the liquid as a solvent, and clear signs of halide-transfer have been found for some compounds with but slight or poorly authenticated solvent properties. It is appropriate, nevertheless, that the aspects of acid-base and solvent properties should be considered in the same context. In the follow­ ing survey, the interhalogen compounds are treated in turn according to formula type, viz. XY, XY3, XY 5 and XY7. Compounds of Formula ΑΎ827.830.831,836,839 Of the compounds of this general formula only ICl and IBr have been described as solvents. One special difficulty attends the interpretation of data on the fused compounds: that is their dissociation into free halogens. For ICl the degree of dissociation is 0-4% at 25°C and 1-1% at 100°C; for IBr the correspondingfiguresare 8-8 and 13-4%. Meaningful physical properties of the liquids are thus sparsely documented (see Table 93). Irrespective of the fact that the potential of chlorine monofluoride as a solvent appears not to have been explored seriously, its ampholytic behaviour is clearly implied by the solid complexes formed through the reactions C1 2 F + AF- <

A

(A = BF 3 or AsF5)925

2C1F

2MF

► 2M + C1F 2 (M = alkali metal or NO)838,84i

[C1F4] + [MF 6 ]-? [BrF 4 ] + [Sb 2 Fn]-*; [BrF 4 ] + [S0 3 F]-? [IF4] + [MF 6 ]-*; [IF4] + [S0 3 F]-?

C1F5,MF5 (M = As, Sb or Pt) BrF 5 ,2SbF 5 ;BrF 5 ,S0 3

IF 5 ,MF 5 (M = Sb or Pt); IF 5 ,117S0 3

IF 7 ,BF 3 ;IF 7 ,AsF 5 ; IF 7 ,3SbF 5

IF 3 I2CI6

CIF5 BrF5

IF 5

IF 7

to

M = alkali metal or NH4; R = organic group. * Crystallographic evidence available. * L. Stein, Halogen Chemistry (ed. V. Gutmann), Vol. 1, p. 133, Academic Press (1967). H. Meinert, Z. Chem. 7 (1967) 41. c R. J. Gillespie and M. J. Morton, Quart. Rev. Chem. Soc. 25 (1971) 553. d C. J. Adams, Inorg. Nuclear Chem. Letters, 10 (1974) in press.

BrF3

[IF 6 ] + [BF 4 ]-; [IF 6 ] + [AsF 6 ]-*; [IF 6 ] + [Sb 3 F 16 ]-

[C1F 2 ] + [BF 4 ]-; [C1F2] + [MF 6 ]-* [BrF2] + [AuF 4 ]-; [BrF2] + [MF 6 ]-*; [BrF 2 + ] 2 [MF 6 ]2[IF2] + [MF 6 ][IC12] + [A1C14]-*; EICl2] + [SbCl 6 ]-*

C1F3

IBr

IC1

C1F3,BF3; C1F3,MF5 (M = P, As, Sb or Pt) BrF 3 ,AuF 3 ; BrF3,MF5 (M = Sb, Bi, Nb, T a o r R u ) ; MF 4 ,2BrF 3 (M = Ge, Sn, Pd or Pt) IF 3 ,MF 5 (M = As or Sb) ICl3,AlCl3;ICl3,SbCl5

[Cl 2 n + [BF 4 ]-;[Cl 2 Fl + [AsF 6 ]-

Formulation

tI2Cl] + [AlCl 4 ]-; [I 2 Cl] + [SbCl 6 ]-

BF3,2C1F; AsF5,2ClF

Composition

Complexes formed with halide ion-acceptors

A1C13,2IC1; SbCl5,2ICl; SbCl5,3ICl

BrCl

C1F BrF IF

Interhalogen

M + [IF 4 ]-;NO + [IF 4 ]M + tICl 4 ]-*;[NR 4 ] + [ICl4]-

MF,IF 3 ;NOF,IF 3 MC1,IC13;[NR4]C1,IC13

M+tIF6]-;NO+tIF6]-; [NR4] + [IF 6 ]M + [IF 8 ]-

MF,IF 5 ;NOF,IF 5 ; [NR4]F,IF5 MF,IF 7 (M=NOorCs) d

M + [BrF 6 ]-

M + [BrF 4 ]-*; Ag + [BrF 4 ]~; NO + [BrF 4 ]"; Ba2 + [BrF 4 "] 2

MF,BrF 3 ;AgF,BrF 3 ; NOF,BrF 3 ;BaF 2 ,2BrF 3

MF,BrF 5

M + [ClF 4 ]-;NO + [ClF 4 ]-

M+[LBr2]"*

MF,ClF 3 ;NOF,ClF 3

ΜΒΓ,ΙΒΓ

M + [BrCl 2 ]-; [NR4] + [BrCl2]M + [IC1 2 ]-*

[NEt4] + [IF 2 ]-

[NEt4]F,IF MCl,BrCl; [NR4]Cl,BrCl (R = Me or Et) MC1,IC1

M + [ClF 2 ]-;NO + [ClF 2 ]-

Formulation

MF,C1F; NOF,ClF

Composition

Complexes formed with ionic halides

TABLE 94. SOLID COMPLEXES FORMED BY INTERHALOGEN COMPOUNDS: (A) WITH HALIDE ION-ACCEPTORS AND (B) WITH IONIC HALIDES*~ C

INTERHALOGEN COMPOUNDS

1519

+

the identity of the ions C12F and C1F2~ having been established primarily by their vibrational spectra. Such behaviour suggests that the meagre electrical conductivity of liquid chlorine monofluoride may arise from the auto-ionization 3C1F^-C1 2 F + +C1F 2 -

though the CI2F + ion appears to be unstable in solution, disproportionating completely, for example, in an antimony pentafluoride-hydrogen fluoride medium even at — 76°C925. 2C1 2 F + -*C1F 2 + +C1 3 +

The slight but significant conductivities of the liquid iodine monohalides have commonly been attributed to auto-ionization schemes of the type 2IX^--I + +IX 2 resembling that suggested for liquid iodine. Altogether more plausible, however, are schemes such as 3IX^— I2X++IX2or 4IX^-I 2 X + +I 2 X 3 -8 40 whereby an interhalogen cation, rather than an isolated monatomic cation, is produced, though it is possible that cations of the type I2X + are extensively disproportionated to give the species I3 + and IX2 + 925. If the auto-ionization is correctly represented in this way, a halide ion-acceptor, which promotes the formation of the interhalogen cation, can be regarded as an acid, whereas a source of halide ions behaves as a base in the liquid iodine monohalide. In that potential acids include the appropriate aluminium(III), tin(IV) or antimony(V) halides, it is noteworthy that phase studies of the systems AICI3-ICI and SbCl5-ICl reveal the formation of the adducts A1C13,2IC1, SbCl5,2ICl and SbCl5,3ICl. A reasonable formulation of these involves the I2C1+ ion, in keeping with the structures of the materials AlC^IClß and SbCl5,ICl3, which, it is known, are better represented as IC12+A1C14~ and ICl2+SbCl6~ (see below). Stable acids belonging to the IBr solvent system have not yet been isolated. On the other hand, the existence of the anions IC12~ and IBr2", for example in crystalline alkali-metal salts and in PCl4+ICl2~, is well established by X-ray diffraction and other measurements; complexes of the type MX,2IX may well contain pentahalide anions I2X3~, as in the case where M = K and X = Cl840. The electrochemical and cryoscopic properties of IC1, IBr and their binary mixtures with organic and inorganic liquids, investigated in detail by Fialkov and his collaborators929»930, are also consistent with the postulated auto-ionization of IC1 and IBr. Thus, the mono­ halides give conducting solutions when dissolved in arsenic trichloride, sulphur dioxide, sulphuryl chloride, nitrobenzene, glacial acetic acid or diethyl ether827»830»831, and it has been shown that, on electrolysis of IC1, for example, in nitrobenzene or acetic acid, both iodine and chlorine appear at the anode in amounts commensurate with the discharge of the IC12 ~ anion. Iodine(I) chloride was first shown to be an ionizing solvent by Cornog and Karges931, who found that the electrical conductivity of the pure liquid is increased considerably when 929 Ya. A. Fialkov and K. Ya. Kaganskaya, / . Gen. Chem. (U.S.S.R.) 18 (1948) 289. wo Ya. A. Fialkov and I. D. Muzyka, / . Gen. Chem. (U.S.S.R.) 18 (1948) 802, 1205; ibid. 19 (1949) 1416; ibid. 20 (1950) 385; Ya. A. Fialkov and O. I. Shor, ibid. 19 (1949) 1787. 931 J. Cornog and R. A. Karges, / . Amer. Chem. Soc. 54 (1932) 1882; J. Cornog, R. A. Karges and H. W. Horrabin, Proc. Iowa Acad. Sei. 39 (1932) 159.

C.I.C. VOL II—AAA

1520

CHLORINE, BROMINE, IODINE AND ASTATINE! A. J. DOWNS AND C. J. ADAMS

potassium or ammonium chloride is dissolved in it; the equivalent conductance of the solute increases with progressive dilution, the values at infinite dilution at 35°C being about 32 and 26 ohm _ 1 cm2, respectively. Other compounds which are appreciably soluble to yield con­ ducting solutions include RbCl, CsCl, KBr, KI, A1C13, AlBr3, PC15, SbCl5, pyridine, acetamide and benzamide. Compounds that are but slightly soluble include LiCl, NaCl, AgCl and BaCl2, while the molecular halides S1CI4, T1CI4, NbCl 5 and SOCl2 dissolve but have little or no effect on the conductivity. The interpretation of the conductivity data is complicated by changes in density, viscosity and degree of dissociation which occur on dilution. Acid-base reactions have been monitored by conductimetric titration. Thus, the titration of RbCl with SbCl5 or of KCl with NbCls shows a break corresponding to 1:1 molar proportions, whilst that of NH4C1 with SnC^ shows a break corresponding to a molar ratio of 2:1, which implies the formation of the ion SnClö 2- . The compound PC15,IC1, which in the solid state is PC14 +IC12 ~, behaves as an acid towards KIC12 and as a base towards SbCls. From the preparative viewpoint, reactions in fused iodine(I) chloride are limited by the prevalence of solvolysis and by the frequent necessity to isolate reaction products by extraction techniques rather than by precipitation; accordingly, pure products are seldom isolable. Nevertheless, it is possible that the molten reagent may prove to be useful in stabilizing chloride complexes of elements in high oxidation states. Fused iodine(I) bromide as a solvent closely resembles the chloride. Alkali-metal bro­ mides furnish polyhalides of formula MIBr2, and phosphorus pentabromide gives the com­ pound PBr5,IBr, which, by analogy with the chloro- compound, is probably PBr4 +IBr2 ~; in IBr this behaves exclusively as a base. Tin(IV) bromide acts as an acid, which may be shown by conductimetric titration to undergo neutralization reactions such as 2IBr 2 ~ + SnBr 4 -> SnBr 6 2 " + 2IBr

Compounds of Formula Ζ73827,830,83ΐ,836,838,839 In this category, studies of solvent behaviour have been restricted to the compounds chlorine trifluoride, bromine trifluoride and iodine trichloride. Physical properties relevant to the solvent functions of the liquids, e.g. dielectric constant, specific conductivity, Trouton constant and viscosity, are contained in Table 93. Chlorine trifluoride is characterized by a very low electrical conductivity and dielectric constant; in general, the signs do not augur well for its capacity as an ionizing solvent. Although no studies of ionic reactions in the liquid have been reported, it is noteworthy that derivatives recently prepared and characterized contain severally the ions C1F2 + 838,925 and CIF4- 838,84i? which, together, are the most likely outcome of the auto-ionization reaction 2C1F3 ^— C1F2++C1F4Thus, compounds of the type MC1F4 (M = alkali metal) are produced, not only by the action of fluorine on alkali-metal chlorides at elevated temperatures, but also by the direct interaction of chlorine trifluoride with alkali-metal fluorides. That these solid complexes contain more-or-less discrete C1F 4 _ anions is implied by their vibrational spectra932. Chlorine trifluoride also forms 1 : 1 complexes with strong Lewis acids like BF3, AsF 5 and SbF 5 , which, on the premises of their vibrational spectra, are best formulated as salts of the C1F2 + cation, e.g. C1F2 + ASFG ", a conclusion recently confirmed by the experimentally 932 K. O. Christe and W. Sawodny, Z. anorg. Chem. 374 (1970) 306.

INTERHALOGEN COMPOUNDS

1521

determined crystal structure of ClF 3 ,SbF 5 (Fig. 28) 933 . Thermodynamic properties have been estimated for some of these complexes on the basis of their dissociation pressures838; for C1F 2 + BF 4 - and C1F 2 + PF 6 -, for example, A77°dissoc = 23-6 and 16-4 kcal mol-i, re­ spectively. The existence of the ion C1F2 + in the liquid phase is also denoted by measure­ ments of the Raman and infrared spectra and electrical conductivities of the liquid mixtures CIF3/HF and ClF 3 /BrF 3 ; thefluorideion-transfer process ClF3 + BrF3 ^ C1F2+ + BrF4" with an equilibrium constant of ca. 10 ~4 at 25°C is thereby implied934. Evidently, therefore, chlorine trifluoride is a significantly weaker fluoride ion-acceptor (i.e. acid) than bromine trifluoride. The 19 F nmr spectrum of the C1F2+ ion in hydrogen fluoride solution has also been described5415. Of the interhalogens, bromine trifluoride is probably the most widely used as a labora­ tory reagent, chiefly for the preparation of complex fluorides. The compound fluorinates everything which dissolves in it, and a discussion of solubilities is therefore restricted to inorganic fluorides. These fall into two groups: I. Alkali-metal fluorides, silver(I) fluoride and barium fluoride; II. Fluorides of gold(III), boron, titanium(IV), silicon, germanium(IV), vanadium(V), niobium(V), tantalum(V), phosphorus(V), arsenic(V), antimony(V), platinum(IV), ruthenium(V) and a few other metals836»838. Members of both groups have been shown to enhance the electrical conductivity of bromine trifluoride, and many form thermally stable adducts with the solvent. To judge by their vibrational spectra934 ~936 and conductivi­ ties in liquid bromine trifluoride, solid complexes of the type MF,BrF 3 (M = alkali metal, Ag, |Ba or NO) derived from fluorides of category I are most aptly formulated as M +BrF4 ~; the presence of a square-planar BrF 4 - anion in the tetragonal crystals formed by KBrF 4 has ultimately been confirmed by neutron-diffraction studies937. By contrast, solid com­ plexes derived from the interaction of bromine trifluoride with the Lewis acids forming category II are best represented as BrF2 + AF - (e.g. A = AuF 3 , AsF 5 , SbF 5 , NbF 5 or RuF 5 ) or [BrF2 + ] 2 AF 2 2 - (e.g. A = GeF 4 , SnF 4 or PtF 4 ). Again, these formulations find support in the infrared and Raman spectra of the solids934»936 and in the crystal structure established for the adduct BrF 3 ,SbF 5 (Fig. 28)938. According to the measured dissociation pressures, the enthalpies of the reactions KBrF4(s) -> KF(s)+BrF 3 (g)

and

BrF2SbF6(l) -> SbF 5 (g)+BrF 3 (g)

are +4-1 and +27-8 kcal mol , respectively838. That the ions BrF 2 + and BrF 4 ~ also exist in the liquid mixtures bromine trifluoride-hydrogen fluoride and bromine trifluoridechlorine trifluoride is intimated, moreover, by the vibrational spectra and electrical con­ ductivities of the liquids934. Accordingly, the consensus of the experimental evidence is that the relatively high specific conductivity of pure liquid bromine trifluoride (> 8-01 x 10 ~3 ohm-i cm-i at 25°C)939 is attributable to the self-ionization _1

2BrF3 ^ — B r F 2 + + B r F 4 933 A. J. Edwards and R. J. C. Sills, / . Chem. Soc. (A) (1970) 2697. 934 T. Surles, H. H. Hyman, L. A. Quarterman and A. I. Popov, Inorg. Chem. 9 (1970) 2726; ibid. 10 (1971) 611,913; T. Surles, L. A. Quarterman and H. H. Hyman,/. Fluorine Chem. 3 (1973/74) 293. 935 j . Shamir and I. Yaroslavsky, IsraelJ. Chem. 7 (1969) 495. 936 K. O. Christe and C. J. Schack, Inorg. Chem. 9 (1970) 1852, 2296. 937 A. J. Edwards and G. R. Jones, / . Chem. Soc. (A) (1969) 1936. 938 A . J. Edwards and G. R. Jones, / . Chem. Soc. (A) (1969) 1467. 939 H. H. Hyman, T. Surles, L. A. Quarterman and A. Popov, / . Phys. Chem. 74 (1970) 2038.

1522

CHLORINE, BROMINE, IODINE AND ASTATINE: A. J. DOWNS AND C. J. ADAMS

It was first pointed out by Woolf and Emeleus940 that dissociation according to this equation implies the possibility of acid-base neutralization reactions in liquid bromine trifluoride. The fact that any fluoride in class I reacts with a member of class II to yield a complex halide is then easily understood in terms, first, of the formation of derivatives of the BrF 2 + and BrF4 ~ ions, and, subsequently, of the interaction of these according to an equation such as B r F 2 + S b F 6 - + A g + B r F 4 " -> A g + S b F 6 " + 2 B r F 3

In a few instances, such as this one, it has been shown by measuring the conductivities of solutions containing different molar proportions of the acid and base that the conductivity exhibits a well-defined minimum at the equivalence point, which corresponds to a molar ratio of 1 : 1. Similar measurements signify that the complex SnF4,2BrF3 reacts with KBrF4 in accordance with the equation [BrF 2 + ] 2 SnF 6 2- + 2 K + B r F 4 " -> K + 2 SnF 6 2" + 4BrF 3

Such reactions form the basis of an extremely useful method for the preparation of complex fluorides, which may be isolated, following neutralization, by evaporation of the excess solvent. Among a very large number of preparations accomplished by this method, the following suffice to illustrate its scope836»838: Ag+Au

BrP8

> AgBrF 4 +BrF 2 AuF 4 -> AgAuF 4

NOCl+SnF4 N204+Sb203 VF5+LiF Ru+KCl

Brp3

> [NO] 2 SnF 6

BrFa

> [N0 2 ]SbF 6

BrF8

►LiVFe

BrP.

►KRuFö

Often it is necessary only to treat equivalent amounts of the acid- and base-forming elements (or compounds of the elements) with bromine trifluoride, which thus functions both as a fluorinating agent and as a reaction medium; the excess of the interhalogen is subsequently removed in vacuo in order to obtain the complex. However, the evaporation stage is not always easily completed: thus, complications may arise through solvolysis or through incomplete interaction of the acid and base according to the neutralization equation BrF2++BrF4" ^2BrF3

High valence states of elements are commonly stabilized in bromine trifluoride. Thus, although palladium(IV) fluoride has not yet been isolated, salts such as K 2 PdF 6 are readily obtained by treatment of alkali-metal chloropalladates(II) with bromine trifluoride; complexes of chromium(V), e.g. KCrOF 4 , have also been made from the corresponding dichromates under analogous conditions. In a similar vein, bromine trifluoride has recently emerged as a solvent for xenon tetrafluoride941, and the formation of complexes of the latter with the acids PF 5 , AsF 5 and SbF 5 has thus been monitored by conductimetric 940 A. A. Woolf and H. J. Emeteus, / . Chem. Soc. (1949) 2865. 941D. Martin, Compt. rend. 268C (1969) 1145.

INTERHALOGEN COMPOUNDS

1523

titration; the solid XeF 4 ,4SbF 5 is said to precipitate from solutions containing XeF 4 and SbF 5 . Bromine trifluoride has lately found a further use as a medium for reaction calorimetry942.

Despite the fact that iodine trifluoride has virtually eluded physical characterization, it is the parent of certain derivatives which are more amenable to investigation, and which testify to fluoride-transfer processes analogous to those described for bromine and chlorine trifluorides. Thus, complexes of the type MF,IF 3 (M = alkali metal or NO) result from the reaction of MF with the trifluoride at low temperatures, or from the action of iodine pentafluoride or fluorine on MI838»854. Presumably to be formulated as tetrafluoroiodates(III), i.e. M + IF4~, these compounds are typically white solids which are thermally stable but decompose readily in the presence of moisture. The abstraction of a fluoride ion from iodine trifluoride also appears to be realized by the action of AsF 5 or SbF 5 at — 70°C. That the 1 : 1 adduct thus formed from SbF 5 is plausibly represented as a derivative of the IF2 + ion is substantiated by the 19 F nmr spectrum of the solid854. Iodine trichloride in the molten state is appreciably dissociated into the monochloride and chlorine, a circumstance which, with other factors, has militated against investigations of its use as a solvent. The measured electrical conductivity of the molten material may well be due in part to the reaction i2ci6^—ICI2++ICI4-

though there is as yet no consistent evidence concerning the nature of the ions produced under these conditions. Approximately square-planar in form, the ICI4" ion has been identified by crystallographic studies of the complex KIC1 4 ,H 2 0 943 , while the crystal struc­ tures of the complexes IC13,A1C13 and ICl3,SbCl5 (Table 40)944 indicate that these are derivatives of the ICI2 + ion. Hence, in its acid-base characteristics, if not in its thermal stability, iodine trichloride evidently has much in common with the halogen trifluorides. Compounds of Formula AT 5 827,830,831,836,838,839 The dielectric constants, degrees of association and electrical conductivities of the three pure liquids (see, for example, Table 93) clearly mark iodine pentafluoride as having the greatest potential as a solvent; the self-ionization implied by the conductivities, which in­ crease in the sequence C1F5 < BrF 5 < IF 5 , is probably best represented by the equation 2XF 5 ^— XF 4 + +XF 6 though the species C1F6~ has yet to be characterized. With the fluoride ion-acceptors AsF 5 and SbF 5 , though not with BF 3 , chlorine pentafluoride certainly forms 1 : 1 adducts, the vibrational spectra of which, it has been argued, support the formulation [CIF4] + [MF 6 ] ~ 945 . On the other hand, no complex attributable to the formation of the C1F6 ~ ion has yet come to light via the interaction of the pentafluoride with either alkalimetal or nitrosyl fluorides. Of the complexes formed by bromine pentafluoride with the Lewis acids SbF 5 and S0 3 , BrF 5 ,2SbF 5 is plausibly assigned the constitution [BrF 4 ] + [Sb2Fn]- on the basis of its crystal structure as well as its infrared and 19 F nmr 942 G. W. Richards and A. A. Woolf, / . Chem. Soc. (A) (1969) 1072. 943 R. J. Elema, J. L. de Boer and A. Vos, Acta Cryst. 16 (1963) 243. 944 c . G. Vonk and E. H. Wiebenga, Acta Cryst. 12 (1959) 859. 945 K. O. Christe and D . Pilipovich, Inorg. Chem. 8 (1969) 391; D . V. Bantov, B. £ . Dzevitskii, Yu. S. Konstantinov, V. F . Sukhoverkhov and Yu. A. Ustynyuk, Doklady Chem. 180 (1968) 491.

1524

CHLORINE, BROMINE, IODINE AND ASTATINE: A. J. DOWNS AND C. J. ADAMS

spectra946a? while BrF5,S03 is presumably either [BrF4]+ [S0 3 F]~ or molecular FS0 2 0BrF 4 . There is also evidence that the 1 : 1 complexes formed by bromine pentafluoride with MF (M = K, Rb or Cs) contain the BrF6" anion: thus, the vibrational spectra of the rhombohedral crystals testify to the presence of such a unit, the bromine atom apparently occupying a site with Du symmetry9461». The neutralization reaction [BrF 4 ] + [ S b 2 F n ] - + 2 C s + B r F 6 ~ -> 3 B r F 5 + 2 C s + S b F 6 -

is reported to occur in the molten complex BrF5,2SbF5 946a . Iodine pentafluoride forms 1 : 1 adducts with SbF5 and PtFs, while giving with SO3 a constant-boiling mixture of composition IF5,1 -HSC^; several oxides and salts of oxy-acids, e.g. N 0 2 , M0O3 and KIO4, also yield definite adducts, the nature of which is as yet obscure. Preliminary reports of the crystal structure947 and vibrational spectrum935 of IF5,SbF5 sustain the representation [IF4] + [SbF 6 ]-, the cation being structurally analogous to the SF4 molecule; a similar constitution is to be expected for IF5,PtF5, while [IF4] +[SC>3F] ~ is possibly an ingredient of the constant-boiling mixture derived from SO3. Solid derivatives of the type MF,IF5 have been described for a variety of univalent cations (M = alkali metal838»839, Ag838»839, NO838»839 or tetra-alkylammonium948). Although no definitive structural analysis has yet been reported, the vibrational949 and 129I Mössbauer950 spectra of the solids support the formulation M + IFÖ ~; the presence of the IFg ~ ion is also implied by the electrical conduc­ tivities and vibrational and 19F nmr spectra of solutions of the salts. In conjunction with a bulky cation such as R4N + (where R = Me or Et) or [(pyridine^X]+ (where X = I, ICI2, Br or Cl), the related anion [IF5C1]~ may well exist in the newly disclosed complexes R4NCUF5 and [(pyridine)2X]Cl,IF5 *5i. A method reported for the preparation of the complex K2PtF6 has exploited bromine pentafluoride as a medium for the reaction of [BrF 2 + ]2[PtF 6 ] 2 - with potassium fluoride, but the solvent properties of chlorine and bromine pentafluoride appear otherwise to have received little attention. On the other hand, the conductivity of liquid iodine pentafluoride is reported to be increased by the dissolution either of a base such as KIF6 or of one of the Lewis acids SbF5, SO3, BF3, HF and KIO3. Neutralization reactions between an acid and base, analogous to those which take place in bromine trifluoride solution, have been established in a few cases, e.g. [IF 4 ] + SbF 6 - + K + I F < r -* K + S b F 6 - + 2IF 5 [IF4]+BF4-+K+IF6- ->K+BF4-+2IF5

Hence, the complex salts KSbF<5, KBF4, KPtFg and CsRhFe have been obtained, but various factors, e.g. incomplete interaction and solvation of the product, render iodine pentafluoride inferior to bromine trifluoride as a reaction medium for the preparation of such complexes. Iodine pentafluoride-hydrogen fluoride mixtures appear to offer signi­ ficantly greater solvent power than pure iodine pentafluoride, and the formation of various complex salts in such mixtures has been described838. According to recent reports, 946(a) M . D . Lind and K. O. Christe, Inorg. Chem. 11 (1972) 608; H . Meinert, U . Gross and A.-R. Grimmer, Z . Chem. 10 (1970) 226; (b) R. Bougon, P. Charpin and J. Soriano, Compt. rend. 272C (1971) 565. 947 H . W. Baird and H . F . Giles, Acta Cryst. A25 (1969) SI 15; N . Bartlett, private communication. 948 H . Meinert and H , Klamm, Z. Chem. 8 (1968) 195. 949 K . O. Christe, J. P. Guertin and W. Sawodny, Inorg. Chem. 7 (1968) 626; S. P. Beaton, D . W. A. Sharp, A. J. Perkins, I. Sheft, H. H . Hyman and K. Christe, ibid. p . 2174; H . Klamm, H . Meinert, P. Reich and K. Witke, Z . Chem. 8 (1968) 393,469. 950 s . Bukshpan, J. Soriano and J. Shamir, Chem. Phys. Letters, 4 (1969) 241. 951H. Klamm and H . Meinert, Z . Chem. 10 (1970) 270.

1525

INTERHALOGEN COMPOUNDS 952

iodine pentafluoride dissolves xenon difluoride and xenon tetrafluoride , and claims have been laid to a number of distinct adducts derived from these systems953. Comparatively weak intermolecular association is clearly implied by the finding that the crystalline com­ pound XeF 2 ,IF 5 contains discrete XeF 2 and IF 5 molecules not perceptibly perturbed by the interaction, which is presumed to be essentially coulombic in origin903. Certain redox reactions have also been reported to occur in liquid iodine pentafluoride. For example, alkali-metal iodides dissolve in the pentafluoride at room temperature with the liberation of elemental iodine and the formation of the tetrafluoroiodate(III) anion, derivatives of which may be isolated from the resulting solution836»838. Iodine dissolves in iodine penta­ fluoride in the presence of a Lewis acid like SbF 5 to form blue solutions until recently supposed to contain the I + cation836»838»839. However, with the characterization of the I 2 + ion925, the absorption spectra and other properties of the blue solutions clearly pro­ claim954 this to be the predominant coloured constituent. Compounds of Formula ΧΥη 836,838,839

Since the triple point is at 6-45°C and the vapour pressure of the solid reaches 760 mm at 4-77°C, iodine heptafluoride is manifestly characterized by a very restricted liquid range. In contrast with previously quoted values, the Trouton constant of the liquid is normal; in view of this and of the meagre values of the dielectric constant and specific conductivity, the liquid offers little promise as a solvent. Of the products of the hypothetical fluoride ion-transfer reaction 2IF7^IF6++IF8-

the IFÖ + cation has been characterized with some assurance. Thus, the heptafluoride has been shown to form the complexes IF 7 ,AsF 5 , IF 7 ,BF 3 and IF 7 ,3SbF 5 , and for the first of these the formulation [IF 6 ] + [AsF 6 ] ~ has been established by spectroscopic950»955»956 and X-ray crystallographic957 measurements. By contrast, there is no sign of interaction of iodine heptafluoride with alkali-metal fluorides, which do not dissolve in the liquid. Al­ though no complexes of the type MF,IF 7 are known, it is noteworthy that IF 7 , IF 7 ,AsF 5 and IF 7 ,3SbF 5 give off fluorine when heated with potassium fluoride to 200-250°C: in these circumstances, neutralization reactions such as [IF 6 ] + [AsF 6 ]-+K + F- ->K + [AsF 6 ]- + IF 7

are believed to precede the decomposition. 3. Halogenating

action

of

interhalogen

compounds827»829 ~ 832 » 838 » 839 .

Except

in

situations where electrophilic attack is encouraged, the halogenating action of an interhalo­ gen compound XYn typically involves the addition or substitution of the more electro­ negative atom Y with simultaneous reduction of the less electronegative atom X, for example to the parent halogen X 2 . The general pattern of behaviour is illustrated by the reactions 952 F . O. Sladky and N . Bartlett, / . Chem. Soc. (A) (1969) 2188. 953 H . Meinert and G. Kauschka, Z. Chem. 9 (1969) 35; V. A. Legasov, V. B. Sokolov and B. B . Chaivanov, Zhur. fiz. Khim. 43 (1969) 2935; A . V. Nikolaev, A. A. Opalovskii, A . S. Nazarov and G . V. Tret'yakov, Doklady Chem. 189 (1969) 982; Doklady Phys. Chem. 191 (1970) 262. 954 R . D . W. Kemmitt, M. Murray, V. M. McRae, R. D . Peacock and M. C. R. Symons, / . Chem. Soc. (A) (1968) 862. 955 K . O. Christe and W. Sawodny, Inorg. Chem. 6 (1967) 1783; K . O. Christe, ibid. 9 (1970) 2801. 956 j . F . Hon and K. O. Christe, / . Chem. Phys. 52 (1970) 1960; M . R. Barr and B. A . Dunell, Canad. J. Chem. 48 (1970) 895. 957 s . P . Beaton, Ph.D. Thesis, University of British Columbia, Vancouver, British Columbia (1966).

1526

CHLORINE, BROMINE, IODINE AND ASTATCNE: A. J. DOWNS AND C. J. ADAMS

of iodine monochloride or monobromide with a wide range of elements E: E+«LX -> EXn+w/2I 2 (X = Cl or Br)

wherein chlorination or bromination, but never iodination, of E is the rule. Likewise the halogen fluorides are noted for their activity as fluorinating agents. Although, in most cases, only qualitative data on the rates and products of the reactions are available, all the halogen fluorides are moderate-to-vigorous fluorinating agents, the approximate order of reactivity being C1F3 > BrF5 > IF 7 > C1F > BrF3 > IF 5 > BrF > IF 3 > IF839: our rather ill-furnished knowledge of the reactions of C1F5 958 suggests that it is somewhat less reactive than CIF3. That such a sequence reflects only in part the thermodynamic potential of the individual fluorides is shown by the molar free energies of reduction listed in Table TABLE 95. MOLAR FREE ENERGIES OF REDUCTION OF THE HALOGEN AND CERTAIN OTHER FLUORIDES

Reaction

AG|98

(kcal mol"1)

ClF5(g)->ClF3(g)+F2(g) IF 7 (g)^IF 5 (l)+F 2 (g) ClF3(g)-+ClF(g)+F2(g) 2ClF(g)-*Cl2(g)+F2(g) BrF5(l)-*BrF3(l)+F2(g) 2BrF(g)->.Br2(l)+F2(g) BrF3(l)->-BrF(g)+F2(g) 2IF(g)->I2(s)+F2(g) IF 5 (l)^IF 3 (g)+F 2 (g) IF 3 (g)^IF(g)+F 2 (g)

+ 9-25» + 14-7* + 15-95» + 27-6a + 27-6» + 35-2» + 39-9* + 56-2» ~ +79*b ~ +80**

XeF6(g)^XeF4(g)+F2(g) XeF4(g)->XeF2(g)+F2(g) XeF2(g)-+Xe(g)+F2(g) 2CoF3(s) -> 2CoF2(s)+F2(g) AsF5(g)-*AsF3(l)+F2(g) 2AgF2(s) -> 2AgF(s)+F2(g) SbF5(l)^SbF3(s)+F2(g) PF5(g)-*PF3(g)+F2(g)

+ 8-1° + 15-2° + 17-9* „ + 44*d + 62-9d·* ~ +69*d ~ +90*b + 161-8d

* Estimated values. ft See Table 93. * L. Stein, Science, 168 (1970) 362. c H. Selig, Halogen Chemistry (ed. V. Gutmann), Vol. 1, p. 403, Academic Press (1967). d National Bureau of Standards Technical Notes 270-3 and 270-4, U.S. Govern­ ment Printing Office, Washington (1968-9). e P. A. G. O'Hare and W. N . Hubbard, / . Phys. Chem. 69 (1965) 4358.

95, which signify that at 25°C the fluorinating capacity should decrease in the sequence ClF5(g) > IF7(g) > ClF3(g) > ClF(g) ~ BrF5(l) > BrF(g) > BrF3(l) > IF(g) > IF5(1) ~ IF3(g). With the changes of phase suffered by the oxidized or reduced forms at different temperatures, some modification of this order is to be expected, but the following general conclusions seem to emerge: (a) in common with the noble-gasfluorides,halogen fluorides 958 D . Pilipovich, W. Maya, E. A. Lawton, H. F. Bauer, D . F. Sheehan, N . N . Ogimachi, R. D . Wilson, F. C. Gunderloy, jun., and V. E. Bedwell, Inorg. Chem. 6 (1967) 1918.

INTERHALOGEN COMPOUNDS

1527

like CIF3 and IF 7 are energetically not far removed asfluorinatingagents from elemental fluorine; and (b) kinetic factors, including the catalytic action of surfaces and of auxiliary Lewis acids like HF, modify the characteristics of thefluorinationreactions to a significant, though as yet largely unspecified, extent. In the gas phase or in solution in a non-polar solvent, homolytic fission of an interhalogen bond is presumably the primary step of halogenation: in a polar solvent, electrophilic attack, either by the electron-deficient site of the undissociated interhalogen molecule or by a cation like CIF2 + , is favoured, particularly with the catalytic assistance of a halide ion-acceptor like HF or C0F3. The intermediate agency of radicals has been substantiated in the decomposition of gaseous chlorine trifluoride861, while several reports testify to the intervention of surface-reactions in processes that purport to occur in the gas phase, but, for the most part, the mechanistic details of such reactions are obscure. No less speculative are the mechanisms of halogenation reactions belonging to the second category, which are accompanied by heterolytic dissociation of an interhalogen bond. In this case, it is the capacity of the interhalogen XYn to act as a donor or acceptor of Y~ ions, rather than the strength of the X-Y bond, which determines the reactivity of the compound. It is likely, though by no means certain, that thefluorinationof elements, oxides and salts of oxy-acids by liquid bromine trifluoride proceeds via fluoride ion-transfer processes of this sort, a view supported by the fluorinating action of derivatives like [BrF2] +[SbF6] - and KBrF4 originating from the interaction of suitable acids or bases with bromine trifluoride839. Cogent evidence of the influence of the environment on the course of halogenation stems from a systematic study of the reaction between iodine monochloride and salicylic acid or phenol959. Thus, the main reaction between the organic compound and the interhalogen vapour involves chlorination since homolytic fission of the I-Cl molecule leads to chlorination by CI2 rather than iodination by the less reactive I 2 ; in carbon tetrachloride solution (low dielectric constant) iodination predominates, accompanied to a small extent by chlori­ nation, heterolytic fission of the I-Cl bond now providing the more facile means of attack; in a solvent of high dielectric constant, such as nitrobenzene, heterolytic fission eclipses the homolytic process, and iodination occurs exclusively. Likewise, with respect to aromatic compounds, iodine monobromide normally acts as a brominating agent as a result of the high degree of thermal dissociation which it experiences and of the relative rates of halo­ genation of the free bromine and iodine thus generated. Only in a solvent of high dielectric constant, like water, does electrophilic iodination occur with this reagent, homolytic dis­ sociation of which can, nevertheless, be suppressed by the addition of elemental iodine839. In general, the interhalogen compounds are less reactive than the parent halogen molecules with respect to homolytic dissociation, but more reactive with respect to heterolytic dis­ sociation. Correspondingly, fluorination reactions which would be difficult to control with elemental fluorine sometimes proceed relatively smoothly when a reagent like chlorine trifluoride, bromine trifluoride or iodine pentafluoride is used as the fluorinating agent. Conversely, the utilization of an interhalogen compound is sometimes advantageous in situations where electrophilic attack by the parent halogen is unduly sluggish. For instance, with respect to electrophilic iodination of organic materials, iodine monochloride is signi­ ficantly more reactive than iodine itself. The most notable applications of the halogenfluoridesstem from their action as fluorinat­ ing agents. Thus, chlorine trifluoride and bromine triflouride have been exploited for the 959 F . W. Bennett and A. G. Sharpe, / . Chem. Soc. (1950) 1383.

1528

CHLORINE, BROMINE, IODINE AND ASTATINE: A. J. DOWNS AND C. J. ADAMS

preparation of binary and complex fluorides—notably of transition metals—from the corresponding elements or their oxides, oxy-salts or halides; reference to some representative reactions has already been made in the context of the solvent action of bromine trifluoride (see p. 1521). Further, as judged by its reactions with sulphur, selenium, molybdenum and tungsten and several of their derivatives928'960-962, chlorine monofluoride has lately emerged as a useful, more moderate fluorinating agent. On the large scale, the halogen fluorides have been widely used in atomic energy installations for the preparation of uranium hexafluoride, and processes have thus been developed for separating uranium from plutonium and fission products in spent nuclear fuels. As attested by a very extensive project literature, chlorine trifluoride and bromine trifluoride appear to be the most serviceable of the fluorides for such recovery processes. Uranium metal and oxides are converted by these agents to volatile uranium hexafluoride, anhydrous hydrogen fluoride being used to accelerate the reaction with chlorine trifluoride. Hence, plutonium is converted to the involatile tetrafluoride, and, tellurium, iodine and molybdenum apart, most fission products also form involatilefluorides,which remain in the reaction vessel when the uranium hexafluoride is volatilized. The uranium hexafluoride is then separated from other volatile components by distillation. The parent halogens and lower-valent halogen fluorides formed by reduction of the trifluoride are treated, in a typical procedure, with fluorine to regenerate the trifluoride. An experimental reactor has even been operated with uranium hexafluoride fuel stabilized with respect to reduction to lower fluorides by the presence of chlorine trifluoride838. Several halogenfluoridesact upon a wide range of oxygen-bearing compounds to liberate oxygen quantitatively, a reaction which has been turned to account for the determination of oxygen in metal oxides and derivatives of oxy-anions, e.g. carbonates, silicates and phosphates. Individual reagents used for this purpose include bromine trifluoride, bromine pentafluoride, KBrF4, [BrF2] +[SbFö] ~ and a gaseous mixture of chlorine trifluoride with hydrogen fluoride838. Extensive use has been made of chlorine trifluoride, bromine trifluoride and iodine pentafluoride for thefluorinationof organic compounds, though the reactions are commonly difficult to control. Moreover, the exceedingly energetic character of many of the reactions of halogen fluorides with hydrogen-containing compounds has prompted investigations of the feasibility of chlorine trifluoride, bromine trifluoride and bromine pentafluoride as oxidizers for such rocket fuels as hydrazine and its derivatives. The system hydrazinechlorine trifluoride has been studied most extensively, since it combines high performance with the advantage over cryogenic systems that both the fuel and oxidizer may be stored as liquids at ambient temperature. The fluorides have also been employed as hypergollic fluids for the ignition of hydrogen-oxygen or solid propellants. Of the three materials so tested, chlorine trifluoride clearly emerges as the most effective, but no development beyond the experimental stage has yet been reported838. Reactions with Elements The reactivity of the interhalogen compounds is indicated by the facts that few elements 960 j . j . Pitts and A. W. Jache, Inorg. Chem. 7 (1968) 1661. 961 C. J. Schack and R. D . Wilson, Inorg. Chem. 9 (1970) 311. 962 R . Veyre, M. Quenault and C. Eyraud, Compt. rend. 268C (1969) 1480; C. J. Schack and W. Maya, / . Amer. Chem. Soc. 91 (1969) 2902; D . E. Gould, L. R. Anderson, D . E. Young and W. B. Fox, ibid. p. 1310; C. J. Schack, R. D . Wilson, J. S. Muirhead and S. N . Cohz, ibid. p. 2907; D . E. Young, L. R. Anderson and W. B. Fox, Inorg. Chem. 9 (1970) 2602.

INTERHALOGEN COMPOUNDS

1529

withstand their action at elevated temperatures, and that many elements react violently, sometimes with spontaneous ignition or explosion, even at or below ambient temperatures. Details of such reactions are to be found in more comprehensive accounts of the compounds826·"829»832'838»839. The rate of reaction is a function, not only of the nature of the reagents and of the temperature, but also of the physical states of the reagents and products, of various catalytic agencies (e.g. solid surfaces, foreign materials or, possibly, one or more of the products of reaction), and, where a solvent is employed, of the nature of that solvent and of the solubilities therein of all the species participating in the reaction. Reaction is likely to be inhibited by the formation of a coherent coating of halide on the surface of a solid element which is either involatile or insoluble in whatever medium surrounds it. Thus, the resistance of nickel and copper and their alloys to attack by the halogenfluoridesdepends upon the protection afforded by such a superficial coating of the corresponding fluoride. In the interaction of an interhalogen compound with an element E exhibiting more than one valence state, oxidation of E commonly proceeds, in conjunction with the halogen of lower atomic number, to the highest valence state, whereas the halogen of higher atomic number is reduced either to the element or to a lower-valent interhalogen derivative. The oxidizing capacity of the interhalogens varies substantially from compound to compound. At one extreme, we find in chlorine trifluoride, for example, an oxidizing agent of unusual power: with but few exceptions, it reacts vigorously with virtually every element at room or elevated temperatures; as in the corresponding reaction with elemental fluorine, the highest known valence state of the element is usually realized in the fluoride thereby produced. With hydrogen, potassium, phosphorus, arsenic, antimony, sulphur, selenium, tellurium, powdered molybdenum, tungsten, rhodium, iridium and iron, spon­ taneous ignition occurs; bromine and iodine also inflame in the trifluoride, forming mixtures of bromine and iodine fluorides (cf. Table 95), while even the noble metals platinum, palladium and gold are attacked at elevated temperatures. Much milder, by contrast, is the action of iodine pentafluoride. Although alkali metals, molybdenum, tungsten, phosphorus, arsenic, antimony and boron react energetically at room temperature or when heated, hydrogen, silver, magnesium, copper, iron and chromium are not disposed to react at normal temperatures. The variation in oxidizing power of the halogen fluorides is clearly delineated by the observation that only certain of them react with the heaviest noble gases xenon and radon. Thus, chlorine trifluoride and iodine heptafluoride, but not iodine pentafluoride, oxidize xenon to xenon fluorides838»963, while in hydrogen fluoride solutions microgram quantities of radon are oxidized at temperatures between —195° and 25°C by all the stable halogen fluorides other than iodine pentafluoride964. While remaining highly reactive materials, iodine monochloride and monobromide lie, nevertheless, at the lower end of the scale of oxidizing capacity. Investigations of their behaviour with respect to a large number of elements965 reveal that moderate-to-vigorous reactions are common, each affording the corresponding chloride or bromide of the element in question, but that numerous elements, including boron, carbon, cadmium, lead, zirconium, niobium, molybdenum and tungsten, fail to react even with the molten inter­ halogen compound. Where reaction occurs, the highest valence state of the element is less frequently realized with these reagents: for example, although phosphorus gives the 963 N. Bartlett, private communication. 964 L. Stein, / . Amer. Chem, Soc. 91 (1969) 5396; Yale Scientific Magazine, 44 (1970) 2; Science, 168 (1970) 362. 965 v . Gutmann, Z. anorg. Chem. 264 (1951) 169; Monatsh. 82 (1951) 280.

1530 CHLORINE, BROMINE, IODINE AND ASTATINE: A. J. DOWNS AND C. J. ADAMS

corresponding pentahalide, vanadium is converted by iodine monochloride, not to the tetrachloride, but to the trichloride, which is conveniently prepared in this way827. Reactions with Oxides and Salts ofOxy-acids The range of halogenating power of the interhalogen compounds is also prominent in their reactions with oxides and salts of oxy-acids. Powerfully oxidizing halogen fluorides, such as chlorine trifluoride or bromine trifluoride, bromine pentafluoride and iodine heptafluoride, typically react with these species to form fluoride or oxyfluoride derivatives, together with oxygen or oxyhalogen species such as C10 2 F or OIF 5 . With chlorine trifluoride the reaction of many metal oxides is sufficiently vigorous to promote ignition, while bromine trifluoride, bromine pentafluoride or the complexes [BrF2] + [SbF 6 ] - and KBrF 4 liberate oxygen quantitatively from numerous oxides and oxysalts, e.g.838 Liquid

6Sb 2 0 3 + 32BrF3 3MRe0 4 +4BrF 3

► 12[BrF2][SbF6] + 10Br 2 +9O 2 Liquid

> 3MRe02F4+2Br2+302

(M = K, Rb, Cs, Ag, £Ca, £Sr or £Ba) KAlSi 3 0 8 + 8BrF5

450°C

► K F + A l F 3 + 3SiF 4 +40 2 +8BrF 3

On the other hand, milder reactions are the rule with iodine pentafluoride: it has been found, for example, that V2O5, Sb 2 0 5 , M0O3, WO3 and C r 0 3 form the following compounds in reactions with the hot or boiling pentafluoride: 2VOF3,3IOF3, SbF 5 ,3I0 2 F, 2Mo0 3 ,3IF 5 , W03,2IF 5 and Cr0 2 F2; while potassium permanganate and perrhenate react thus: KMO4+IF5 -> M 0 3 F + I O F 3 + KF (M = Mn or Re)

Included within the same general category are the reactions of glass and quartz, which occur more or less readily with all the halogen fluorides, e.g. 2B203+4BrF3->4BF3+2Br2+302 3Si0 2 +4BrF 3 - > 3 S i F 4 + 2 B r 2 + 3 0 2 Si02+2IF7

-► SiF 4 +20IF 5

It is possible that attack is usually initiated by traces of hydrogen fluoride and sustained by a continuous cycle of reactions such as 4 H F + S i 0 2 -» S i F 4 + 2 H 2 0 I F 7 + H 2 0 -> OIF5+2HF

This mechanism is consistent with the observations that very pure samples of chlorine trifluoride or iodine pentafluoride are slow to etch glass or quartz apparatus. Again, the action of hydroxyl compounds illustrates well the diversity of oxidizing proper­ ties of the interhalogen compounds966. Strongly oxidizing halogen fluorides like chlorine trifluoride interact violently with water unless precautions are taken to moderate the process, for example by the use of low temperatures and by dilution; the products formed depend partly upon the proportions of the reactants, but may include HF, HX, HOX, HXO3, 0 2 and X 2 , as well as halogen oxides (e.g. C102 and OF2) and oxyhalogen fluorides (e.g. C102F) in certain cases. Whereas iodine pentafluoride also suffers vigorous hydrolysis, the iodine

INTERHALOGEN COMPOUNDS

1531

does not undergo reduction: 3 H 2 0 + I F 5 -> H I 0 3 + 5HF

ΔΗ2η = - 2 2 0 5 kcal mol"*

Altogether more moderate are the hydrolytic reactions of the diatomic interhalogens BrCl, ICl and IBr and of iodine trichloride. The primary steps in the hydrolysis of the diatomic molecules in the absence of added halide ions are: and

-> H + + X - +ΗΟΙ (X = Cl or Br)

IX+H20

BrCl+H 2 0 -> H + +C1- +HOBr

The reactions occurring during the hydrolysis of iodine trichloride are complicated, but have been formally represented by the equation 2IC1 3 +3H 2 0 -+ IO3- +IC1 2 - +4C1- + 6H +

However, in acidic solutions containing excess of halide ions, hydrolysis of the interhalogen yields place to the formation of relatively stable polyhalide ions such as ICI2 ~ and IBr2 ~. In the form of the IC12 ~ ion, present in strongly acid solution, iodine monochloride features in volumetric analysis, being for example the reduction product of iodate in the well-known Andrews titration. Iodine trichloride gives as possible products Cl~, ICl4 ~, ICl2 ~ and 10 3 ~ in proportions depending on the precise conditions of reaction839. Although Br-Cl, I-C1 and I-Br bonds can thus survive hydrolysis, halogen-fluorine bonds appear invariably to suffer hydrolysis, this despite the enhanced energy of such bonds. Under aqueous condi­ tions, therefore, an electropositive halogen atom has the characteristics of a class " b " acceptor328, though, as noted elsewhere (Section 3, p. 1252), this property hinges upon differences of bond energy and of hydration energy, being related only indirectly to intrinsic properties of the interhalogen bonds. With certain compounds containing multiple bonds between carbon and oxygen or between sulphur and oxygen, the diatomic molecules C1F and BrF undergo addition reac­ tions. In some instances, the addition is localized on the carbon or sulphur atom, e.g. co +

CIF

FC1CO

SO, -f

XF

X O S F O

m

(X^=C1 or Br )

but in others addition occurs across one of the multiple bonds, e.g. Ri

\ 0=0+CIF /

-> RjRzCFOCl

R2

(Ri, R2 = F, CF3 or other perhalogenoalkyl groups)962 F 4 S = 0 + CIF -> SF5OC1962 SO3 + CIF

-> FS02OC196i

See for example K. O. Christe, Inorg. Chem. 11 (1972) 1220.

1532 CHLORINE, BROMINE, IODINE AND ASTATINE: A. J. DOWNS AND C. J. ADAMS

Reactions of the latter class provide an expeditious route to molecular hypochlorites, being catalysed either by a base like caesium fluoride or by an acid like hydrogen fluoride962. Reactions with Halides Fluorination, frequently allied to oxidation, is innate to the reactions of the halogen fluorides with many other halides derived from metals or non-metals. For example, the chlorides NiCl2, AgCl and CoCl2 are converted to NiF 2 , AgF 2 and C0F3 by chlorine trifluoride at 250°C838. Arsenic trichloride is oxidized by chlorine mono- or trifluoride to the compound formulated as [ASCI4]+[AsF6] ~, while with antimony pentachloride, chlorine monofluoride gives SbCl4F. The addition of chlorine monofluoride to sulphur tetrafluoride at 350°C affords one of the most efficient methods yet devised for the prepara­ tion of pentafluorosulphur chloride928; at 180°C chlorine trifluoride converts sulphur tetrafluoride to a mixture of pentafluorosulphur chloride and sulphur hexafluoride838. Liquid bromine trifluoride acts upon many metallic chlorides, bromides and iodides to produce binary or complex fluorides in which the highest accessible valence state of the metal is often attained, though the formation of insoluble surface films of fluoride may prevent the reaction from going to completion, as with lead(II), thallium(I) and cobalt(II) halides, which yield mixtures of lower- and higher-valent fluorides. However, when the product is volatile (e.g. UF 6 ) or soluble in the reagent (e.g. the fluorides of the alkali metals), conversion into the fluoride is quantitative. Reactions with Hydrogen-containing and Organic Compounds Many hydrogen-containing compounds, both organic and inorganic, inflame or explode when mixed with a halogen fluoride as a result of highly energetic reactions such as 838 2NH 3 + 2C1F3 - > N 2 + 6HF+C1 2 and 3N 2 H 4 +4C1F 3 -> 3N 2 +12HF+2C1 2

On the other hand, some of the reactions may be moderated by diluting the interhalogen with an inert gas, by dissolving the reagent in a relatively inert solvent such as carbon tetrachloride or a fluorocarbon, or by lowering the temperature. For example, all the chlorine fluorides convert difluoroamine to chlorodifluoroamine, e.g. CIF3 + 3HNF 2 -> C1NF 2 +N 2 F 4 +3HF

in reactions which can be moderated by the use of complexes of either the chlorine fluoride or of the difluoroamine967. Chlorodifluoroamine is also the product of hazardous reactions involving chlorine trifluoride and ammonium fluoride or bifluoride838. The reactions of the halogen fluorides with organic compounds have been reviewed by Musgrave926, and will be described here only in outline. Even where moderating conditions are used, substitution reactions in which the chlorine, bromine or iodine of organo-halogen compounds is replaced by fluorine are generally less exothermic and easier to control than reactions involving C-H bonds. However, cleavage of C-C bonds can also occur, particu­ larly with the more energetic reagents, and violent explosions have been reported as a result of contact between chlorine trifluoride and relatively inert hydrocarbons, e.g. Kel-F oil838. Even with iodine pentafluoride, the mildest of the halogen fluorides, compounds rich in hydrogen inflame at room temperature. Nevertheless, numerous organo-halogen 967 D. Pilipovich and C. J. Schack, Inorg. Chem. 7 (1968) 386.

INTERHALOGEN COMPOUNDS

1533

compounds undergo smooth substitution reactions with iodine pentafluoride: for example, CCI4 is slowly converted to CCI3F together with some CC12F2, CHI 3 gives mainly CHF 3 , while CF3I is the principal product of the reaction with CI 4 830 . Other organic compounds whose reactions with one or more of the halogen fluorides have been investigated include carbon tetrabromide, carbon disulphide, trichloroacetic acid, polychlorobutadienes, acetonitrile and related compounds, aliphatic amines, various ketones, as well as benzene, hexachlorobenzene and other aromatic compounds827»830'838»926. Substitution reactions are the general rule, but in some instances, as in the vapour-phase fluorination of benzene and toluene by chlorine trifluoride, both addition- and substitution-products are formed. Many of the reactions have been the subjects of patents; certain of them, e.g. between iodine pentafluoride and carbon tetraiodide or carbon disulphide, played an important part in providing, for the first time, routes to perfluoroalkyl derivatives926, but, in the main, these have now been supplanted by procedures more amenable to control and less severe in their technical demands. As noted previously, the course of substitution reactions involving the compounds BrCl, IC1 and IBr varies with the nature of the medium employed. Most extensively studied has been the electrophilic halogenation of C-H bonds in aromatic systems. In particular, the reactions of iodine monochloride with, inter alia, the following compounds have been studied in some detail: phenols, aromatic carboxylic acids, aromatic ethers, acetanilide and other N- or ring-substituted aniline derivatives, aromatic amino-acids and amides, and various thiazole derivatives. Iodine trichloride also reacts under appropriate conditions with various aromatic compounds, including thiophen, to give chloro-substituted products, usually together with no more than small amounts of the corresponding iodo-substituted derivatives. With certain organometallic derivatives, substitution at the metal-carbon bond may take a somewhat different course. Thus, iodine monochloride is reported to act upon organo-mercury compounds in various solvents to give organo-iodine, rather than organochlorine, compounds, unless secondary reactions intervene, whereas iodine trichloride converts aryl-tin or aryl-mercury compounds into the corresponding diaryliodonium derivatives (q.v.), e.g. ICl 3 + 2PhSnCl3 -> Ph 2 ICl+2SnCl 4

A characteristic reaction of the diatomic interhalogen molecules is addition to the units ^;C=C.

or

-0Ξ=Ο-

of organic compounds. Interaction of this kind has been reported,

not only for the compounds C1F, BrCl, IC1 and IBr, but also for the less well characterized fluorides BrF and IF. As produced in situ by the reaction of the parent halogen either with silver(I) fluoride968 or with a higher-valent halogen fluoride, e.g. BrF 3 or IF5926»927, the monofluorides add to the double bonds of various olefins926 9 2 7 , 9 6 8 or of unsaturated carbohydrate derivatives like D-glucal triacetate968. Numerous addition reactions with olefinic and acetylenic compounds have also been reported for bromine and iodine mono­ chloride; indeed the products of such reactions afforded some of the earliest positive evidence for the existence of bromine monochloride as a distinct compound. The reactions with iodine monochloride have been successfully exploited to determine the degree of unsaturation (the "iodine number") in natural and synthetic rubbers. The orientation of addition and the kinetic characteristics of such reactions are consistent with a mechanism 968 L. D. Hall and J. F. Manville, Canad. J. Chem. 47 (1969) 361, 379; L. D. Hall and D. L. Jones, ibid, 51 (1973) 2902.

1534 CHLORINE, BROMINE, IODINE AND ASTATINE: A. J. DOWNS AND C. J. ADAMS

analogous to that proposed for the addition of a homonuclear halogen molecule (see p. 1218), electrophilic attack being initiated by the more electropositive halogen atom which forms a cationic organo-halogen intermediate. Despite the instability of the pure compounds, the propensity of BrN3, IN 3 and INCO to add to olefinic species has also been established, the pseudohalide molecule typically being generated in situ by the reaction of the parent halogen with a salt of the pseudohalide anion874; it is reported that, with appropriately chosen conditions, BrN3 can undergo either electrophilic or free-radical addition, with correspond­ ing variations in the orientation of the bromine and azide substituents, but that IN 3 acts exclusively as an electrophile. Iodine trichloride adds to acetylene apparently to give the compound ClCH=CHICl2, and also to other hydrocarbons containing one or two isolated double bonds or two conjugated double bonds, though the nature of the products is not well defined827. 3. P O L Y H A L I D E ANIONS826-829,832,834,835,840,841,857

Introduction and Classification A polyhalide anion may be defined as an addition product of a halide ion acting as a Lewis base with one or more halogen or interhalogen molecules acting as Lewis acids. Such an ion can be assigned the generalized formula XiVynZv ~, where X, Y and Z represent either identical or different halogen atoms. For all the well-defined species of this type, m +n +p is an odd number that can be 3, 5, 7 or 9. Hence, it is clear that the anion contains an even number of electrons. The apparent anomaly posed by the complex " C s V is resolved by the discovery that the material is diamagnetic and contains the Ig2~ ion. Other complex species have been alluded to in which the sum m+n+p appears either to exceed 9 or to be an even number, but, with very few exceptions, convincing evidence of the formulation or structure is lacking. TABLE 96. POLYHALIDE ANIONS XmY„Zp ~ FOUND IN CRYSTALLINE SALTS*

m+n+p= 3 I3I2Br~ I2C1IBr2" ici2IBrCl" IBrF" IF2-* Br3" Br2CI" BrCl2" BrF2"c CI3C1F2-

5 I5I4CII 4 BrI 2 Br 3 I 2 Br 2 ClI 2 BrCl 2 I2CI3IBrCV ICI4ICI3FIF4BrF 4 " CIF4-

Other species 7 I7I 6 BrIF 6 Br 6 ClBrF 6 "

9 I9IF 8 " d

Is 2 "

a A. I. Popov, Halogen Chemistry (ed. V. Gutmann), Vol. 1, p. 225, Academic Press (1967). b H. Meinert, Z. Chem. 7 (1967) 41. c T. Surles, L. A. Quarterman and H. H. Hyman,./. Inorg. Nuclear Chem. 35 (1973) 668. d C. J. Adams, Inorg. Nuclear Chem. Letters, 10 (1974) in press.

POLYHALIDE ANIONS

1535

Table 96 lists the polyhalide anions identified more or less convincingly in the form of crystalline complexes; the catalogue includes neither the anions of compounds whose identification remains dubious nor those solution species whose identification rests solely on physicochemical measurements of stability constants. Crystalline derivatives of the anions, as of the hydrogen dihalide anions (Section 3, p. 1313), are formed most easily with large univalent cations, e.g. Cs + , NR4 + (R = alkyl group) or ASPI14+. Many crystalline polyhalides also contain solvent or other foreign donor molecules, which are apparently essential to the stability of the compound, since the material suffers irreversible decomposi­ tion when attempts are made to remove the molecules. Examples of these solvated poly­ halides include KI 3 ,H 2 0, KI 7 ,H 2 0, KIC1 4 ,H 2 0, MI3,2PhCN (M = Na or K) and Kl3,2MeHNCHO. Such complexes should not be confused with the polyhalide salts formed by many organic bases which suffer either protonation or coordination to a cationic halogen-centre to accommodate the formation and stabilization of the polyhalide anion, e.g. protonation

► [2,2'-bipyridylH] + ICl 2 -

2,2'-bipyridyl,2ICl

2C5H5N,IC1

polar solvent

> [(C 5 H 5 N) 2 I] + [IC12]"

Preparation The general method for the preparation of polyhalides simply involves direct combina­ tion between an ionic halide and the appropriate halogen or interhalogen, the experimental conditions being dictated by the reactivity of the reactants and products. This may be brought about (i) by exposing the solid halide to the gaseous or liquid halogen or interhalo­ gen; (ii) by mixing the parent halogen or interhalogen compound with a halide or polyhalide in a suitable solvent, e.g. water, an alcohol or acetonitrile; (iii) by producing the interhalogen compound and effecting combination in situ; (iv) by displacement of a halogen of higher atomic number in an existing polyhalide by one of lower atomic number. These general methods have been comprehensively described elsewhere840»841 »857,969. Similar methods have been used to prepare pseudohalide analogues of the polyhalide ions: for example, oxidation of solutions containing iodine and thiocyanate ions gives rise to the anion [I(SCN)2] ~970. In various circumstances, polyhalide ions may well have an intermediate agency; thus, there are spectroscopic reasons for believing that, when silver iodide crystals are subjected to high pressures, 13 _ ions are formed, presumably as a stage in the evolution of metallic properties971. The gas-solid reactions are, for the most part, comparatively slow and inefficient, often failing to produce pure products. On the other hand, many polyhalide complexes (excluding those containing fluorine) can be successfully prepared in methanol or ethanol solutions. With the more reactive polyhalides, a solvent less susceptible to halogenation, e.g. 1,2dichloroethane, has to be selected. As previously noted (p. 1515), derivatives of fluorinecontaining polyhalide anions can commonly be gained by the direct action of the liquid halogen fluoride on a suitable ionic fluoride; alternatively, the fluorinating action of the halogen fluoride or of elemental fluorine on chloride, bromide or iodide derivatives is 969 A . I. Popov and R. E . Buckles, Inorganic Syntheses, Vol. 5 (ed. T. Moeller), p. 167, McGraw-Hill, N e w York (1957). 970 c . Long and D . A . Skoog, Inorg. Chem. 5 (1966) 206. 971 M. J. Moore and D . W. Skelly, / . Chem. Phys. 46 (1967) 3676.

1536 CHLORINE, BROMINE, IODINE AND ASTATINE: A. J. DOWNS AND C. J. ADAMS

turned to account in reactions such as 838 » 841 90-250°C

and

MC1+2F2

^MC1F4 (M = K, Rb or Cs)

liquid B r F 3

6MCl + 8BrF3 M5MBrF4+3Cl2 + Br2 (M = K, Ag or £Ba) 857 Contrary to earlier reports , it now appears that not only halogen, but also interhalogen, molecules add to appropriate polyhalide ions, e.g. IC12- +XC1 -> IXCI3- (X = Cl or I) It has thus been possible to prepare a variety of mixed pentahalide salts, e.g. M[I 2 X 3 ] and M[I 2 X 2 Y], where M + = NR 4 +, [C 5 H 5 NH]+, [2,2'-bipyridylH] + or [1,10-phenanthrolineH]+ and X, Y = Cl or Br8*o. It is relatively simple to prepare polyhalide salts of anions like I 3 - , IC1 2 ~ and IBr 2 ~, which are comparatively stable with respect both to dissociation and to halogenation reactions, but the preparative difficulties increase rapidly as the chemical stability of the anion deteriorates. For example, the Cl3~ ion is comparatively unstable with respect to the dissociation CI3- ^ C12 + C1 _ , whether in the solid state or in solution; as a result, salts of this anion are scarce and rather poorly characterized. Some of the salts of fluorinecontaining anions can be formed only at low temperature, and the difficulty of preparation is further aggravated by their vigorous activity as fluorinating agents and by their suscepti­ bility to hydrolysis. In the formation of polyhalide complexes, the final product is usually determined more by the composition of the reaction mixture than by the precise nature of the reagents. Thus, potassium dichloroiodate(I) can be obtained by the reaction (i) of potassium iodide with chlorine, (ii) of potassium chloride with iodine monochloride, or (iii) of potassium tetrachloroiodate(III) with potassium iodide. When a polyhalide ion engages in a replace­ ment reaction, the terminal atoms are typically displaced by more electronegative substituents. For example, the reaction of bromine with the I2Br~ or I 3 - ions invariably leads to the formation of IBr2 ", but reaction does not proceed beyond this stage to give Br3 ~. Likewise the reaction Et4NICl2+2AgF

aoetonltrlle

> Et4NIF2+2AgCl has been reported839. Phase studies of binary systems, such as halogen- or interhalogen-alkali halide, and of ternary systems, such as halogen- or interhalogen-alkali halide-water, have been used by numerous investigators for the identification and characterization of polyhalide salts. In a number of cases the experimental results have clearly established the existence of individual polyhalides, but, on occasion, claims have been made for compounds which have not been substantiated and whose existence or nature still remains open to question84*). The tendency of polyhalides to form solvated solids by interaction with water and other donor molecules evidently becomes more pronounced as the bulk of the cation decreases. Thus, whereas rubidium and caesium readily form solvent-free tri-iodides, the corresponding derivatives of lithium, sodium and potassium are known only as solvated materials, e.g. LiI3,4PhCN and KI 3 ,H 2 0; similarly, solvent-free salts of the IC1 4 ~ anion are formed by all the alkali-metal cations other than lithium. In the formation of solvent-free polyhalides, there is ample experimental evidence that the stability of the solid complex is governed by

POLYHALIDE ANIONS

1537

the nature of the cation. Thus, to judge by the efforts made to prepare such complexes, small or multiply charged cations are generally ill-adapted, while bulky univalent organic cations such as NR 4 + , [C 5 H 5 NH] + , SR3 + and AsPh 4 + are most conducive, to the stabilization of polyhalide anions. This generalization appears to be valid in spite of certain irregularities imposed by the nature of the polyhalide anion: for example, tetrachloroiodate(III) derivatives of the alkaline-earth and divalent transition metals have been prepared, as have Ni(ClF2)2972, Ni(ClF4)2972, Ba(BrF4)2, [Ni(NH3)6][I7]2 and [CO(NH 3 )Ö][I 3 ] 3 . With a suitable cation, crystalline polyhalides containing up to 9 halogen atoms, as in NMe4l9, have been characterized. Curiously, only the I3~, I5~ and I9~ salts can be obtained with the cation NMe 4 + and the I3~, I5~ and I7~ salts with NEt 4 + . However, the complex derived by precipitation or crystallization is determined not only by the forma­ tion constant of the anion, but also by the relative solubilities of the different components, which bear a subtle and largely obscure relation to the properties of the ions. Hence, preparative success is not necessarily a guide to the thermodynamic stabilities of complex polyhalides in the solvent-free condition. Moreover, there are good reasons for believing that bulk and charge are not the only properties of the cation which affect the stability of solid polyhalides; due weight must also be given to the precise shape, polarizability and charge-distribution of the cation, though this has seldom been feasible at more than a qualitative, empirical level in studies up to the present time. The reaction of a halogen or interhalogen compound with a molecular halide sometimes entails the formation of a crystalline addition compound. Depending on the relative power of the reagents to accept halide ions, there are two ways of formulating the product, viz. or

MXn+2XY -+ [XY 2 ] + [MX» + i ] MXn + XY - * [ Μ Χ η - ΐ ] + [Χ 2 ΥΓ

which may thus contain either a polyhalogen cation or a polyhalide anion; in practice, the difference between these two conditions may well be clouded by the formation of a polymeric network of the type found in ClF3,SbF5 (see Fig. 28). The situation is further complicated in certain cases where amphoteric behaviour appears to be possible, as with the complex IC1,PC15, and, depending on experimental conditions, halide ion-transfer can take place in either direction. Formed from phosphorus tri- and pentahalides are complexes such as IX,PC15, IX,PBr5 and Br2,PBr5. That the solids contain trihalide anions has been verified in the cases of IC1,PC15 836,840 a n ( j Br2,PBr5 973 , which are to be formulated, according to crystallographic analysis, as [PC14] +[IC12] ~ and [PBr4] +[Br3] -, respectively. By contrast, the power of antimony(V) halides as halide ion-acceptors appears to favour the formation of cationic polyhalogen derivatives. The protonic acids corresponding to the polyhalide ions have not been prepared in the pure state, and it is doubtful whether the stabilities of the materials are compatible with isolation under conventional conditions. The acids are presumably formed in solutions whenever a halogen or interhalogen is added to a hydrohalic acid; the limited evidence available suggests that their strength as proton-donors is superior to that of the hydrohalic acids. Aqueous solutions containing up to 94% of HIBr2 and 80% of HIBrCl or HIC12 have been prepared, but attempts to isolate crystalline phases containing the acid have been unsuccessful. Orange-yellow crystals of the hydrate HIC14,4H20 result on cooling the 972 L . Stein, J. M . Neil and G. R. Alms, Inorg. Chem. 8 (1969) 1779. 973 G . L . Breneman and R . D . Willett, Ada Cryst. 2 3 (1967) 467.

1538 CHLORINE, BROMINE, IODINE AND ASTATINE: A. J. DOWNS AND C. J. ADAMS

solution formed by passing chlorine through a suspension of iodine in concentrated hydro­ chloric acid; the material decomposes when attempts are made to dehydrate it. Representa­ tive of the solids which have been identified as complexes of a protonic acid with an appropriate nitrogen-base are HI 2 Cl3,C 5 H 5 N, HI 3 ,2PhCONH 2 and HI 3 ,4PhCN; on the evidence available the proton is accommodated either as a discrete cation, e.g. [C5H5NH]+, or within a hydrogen-bonded network of some kind. Spectrophotometric investigation of the hydrogen iodide-iodine system in carbon tetrachloride suggests an equilibrium constant for the reaction HI + I 2 ^ HI 3 in the range 25-400840. Physical Properties The physical characterization of polyhalide species has assumed three principal aims: (i) qualitatively or quantitatively the thermodynamic stability with respect to dissociation has been assessed for many polyhalides either as solids or in solution; (ii) the structures of numerous solid complexes have been determined; and (iii) spectroscopic properties of the complexes in the solid or solution phases have been evaluated. 1. Stability of the Polyhalides Solids A solid polyhalide derivative of the type M +[XwYnZp] ~ tends to dissociate spontane­ ously into a halogen or interhalogen compound and a simple monohalide derivative of the cation M + . In many cases, the dissociation pressure is already appreciable at room tempera­ ture. In accordance with this behaviour, the chemical reactions of the polyhalides are mostly those of the dissociation products. The relative thermal stabilities of crystalline trihalides have been assessed by comparing the relative magnitudes of the dissociation pressure at a given temperature or the temperatures at which the dissociation pressure reaches a particular value (e.g. 1 atm). An alternative approach has consisted of suspending the polyhalide salt in an organic solvent like carbon tetrachloride and determining the equilibrium concentration of free halogen or interhalogen in the liquid phase; since both the poly- and monohalide salts are virtually insoluble in non-polar organic liquids, this concentration is a direct index to the degree of dissociation of the polyhalide. More recently, the free energy change accompanying a reaction of the type Mlx+2y(s) ->

MIz(s)+yl2

(M = Rb, Cs, NH 4 , NMe 4 or NEt 4 ; x = 1, 3 or 5 and y = 1 or 2)

has been determined974 from emf measurements of the solid-state galvanic cell A g | A g I | C , MI*+ MI* +2y Ag|AgI|C,I2(s)

over the temperature range 25-113-5°C; hence, the results presented in Table 97 have been derived. On the basis of the various measurements, the following generalizations may be made: (i) For several series of polyhalides containing different cations and the same anion, thermal stability is promoted by increasing the size of the cation. 974 L . E. Topol, Inorg. Chem. 7 (1968) 451; ibid. 10 (1971) 736.

1539

POLYHALIDE ANIONS TABLE 97. VALUES OF AG°, ΔΗ° AND Δ5° AT 25°C FOR THE REACTION

MII(s)+^I2(s)-*MII+2l((s)''

-AG° -AH° (kcalmol-1) (kcalmol-1)

Δ5° (caldeg-1 mol-i)

Reaction

Cation

MI(s)+I2(s)-*MI3(s)

NH4 Rb Cs NMe4 NEt4

1-8 2-4 3-5 3-5 5-5

21 31 3-7 1-7

-11 -2-3 -0-9 + 6-7

MI3(s)+iI2(s)->MI4(s)

Cs

0-63

0-80

-0-6

MI3(s)+I2(s)->MI5(s)

NMe4

3-21

2-2

+ 3-2

MI3(s)+2I2(s)->MI7(s)

NEt4

5-27

5-5

-0-9

MI5(s)+2I2(s)->MI9(s)

NMe4

2-14

0-3

+ 6-2

a

L. E. Topol, Inorg. Chem. 7 (1968) 451; ibid. 10 (1971) 736.

(ii) On dissociation, the solid monohalide afforded by the smallest halide ion results. Thus, CsICl2 gives CsCl + ICl rather than Csl + Cl2. (iii) For polyhalides containing the same cation but different trihalide ions, the precise stability sequence varies according to the method used to assess the extent of dissociation, but measurements of dissociation pressure imply the following order of decreasing stability: I 3 - > IBr2" > IC1 2 - > I 2 Br- > Br 3 " > BrCl2" > Br2Cl~. Evidently the most stable anions are those containing iodine as the central atom, while the stability of symmetrical anions, e.g. IBr2~, is superior to that of related but unsymmetrical anions, e.g. I2Br~. Solutions The behaviour of the polyhalide ions in solution has been the subject of many investiga­ tions, though in comparatively few cases have reliable thermodynamic properties been established for the solution species. The solvents employed include, in addition to water, chloroalkanes, e.g. C1CH2-CH2C1 or CO^Br, alcohols, alcohol-water mixtures, anhydrous or aqueous acetic acid, trichloroacetic acid, acetonitrile, acetone and nitromethane. Studies of the solutions are complicated by the fact that the properties of polyhalide ions are drasti­ cally influenced by the nature of the solvent; in particular, the possibility of solvolysis or of halogenation of the solvent must be taken into account. Water is a far from perfect medium because of the susceptibility to hydrolysis of the polyhalide anions and of the halogens or interhalogens formed by the dissociation reaction. While hydrolysis can be more-or-less suppressed by the addition of an acid in fairly high concentration, this not only increases the ionic strength, but also introduces another reactive species into the system. The following physicochemical methods have provided the experimental mainstay for the characterization of the polyhalide ions in solution: spectrophotometry (mostly in the ultraviolet and visible regions), electrical conductance, potentiometry, polarographic and voltammetric techniques, liquid-liquid distribution experiments, and measurements of solubilities or colligative properties. Hence, it has been possible to establish the existence

1540 CHLORINE, BROMINE, IODINE AND ASTATINE: A. J. DOWNS AND C. J. ADAMS

in aqueous solutions of trihalide ions corresponding to all the possible combinations of chlorine, bromine and iodine atoms. Even the highly unstable CI3 ~ ion exists in concentrated chloride solutions saturated with chlorine. By contrast, if fluorine-containing polyhalide ions are capable of existing in aqueous media, convincing evidence of the fact has still to be found. In media where the opportunity of solvolysis is denied, e.g. bromine trifluoride934 or acetonitrile949, the presence of fluoro-anions like B r F ^ and IF 6 ~ is clearly intimated by the vibrational spectra or electrical conductivities of the solutions; of the stability of such species with respect to dissociation, however, little is known. TABLE 98. THERMODYNAMIC CONSTANTS FOR THE FORMATION OF POLYHALIDE ANIONS IN AQUEOUS SOLUTION* " e

Ion

Formed from

Formation constant at 298°K(Ä,) 710* 480* 166» 34 d 16-3* 10-5d 1-66» 114° 019» 7-2xl0-3*c

l3"(aq) IBr2"(aq) ICl 2 -(aq) IBrCl-(aq) Br3"(aq) I2Br~(aq) I 2 Cl-(aq) Br 2 Cl-(aq) Cl 3 -(aq) BrCl 2 _ (aq)

I 2 (aq)+I-(aq) IBr(aq)+Br~(aq) ICl(aq)+Cl-(aq) IBr(aq)+Cl~(aq) Br 2 (aq)+Br"(aq) I 2 (aq)+Br"(aq) I 2 (aq)+Cl-(aq) Br 2 (aq)+Cl-(aq) Cl 2 (aq)+Cl-(aq) Br 2 (aq)+Cl-(aq)

I 2 Cl-(aq) I 2 Br-(aq)

ICl(aq)+I-(aq) IBr(aq)+I"(aq)

3 X 108e 2-6X106d

I 5 -(aq) Br5"(aq)

I 2 (aq)+I 3 "(aq) Br 2 (aq)+Br 3 -(aq)

~9a 1.45a

— AG°298

(kcalmol - 1 ) 3-9 3-7 30 21 1-68 1-4 0-3 008 -1-0 11-6 8-8 1-3 0-2

— Δ / f °298

(kcalmol" 1 )

Δ5°298

(caldeg - 1 mol-i)

4-5

-2-2

1-50 1-4

+0-6 -0-3

0

+0-4

— 2-2

— -70

*K, = [BrCl2-][Br-]/[Br2][Cl-J2. a A. I, Popov, Halogen Chemistry (ed. V. Gutmann), Vol. 1, p. 225, Academic Press (1967). National Bureau of Standards Technical Note 270-3, U.S. Government Printing Office, Washington (1968). c R. P. Bell and M. Pring, / . Chem. Soc. (A) (1966) 1607. d R. Guidelli and F. Pergola, / . Inorg. Nuclear Chem. 31 (1969) 1373. e G. Piccardi and R. Guidelli, / . Phys. Chem. 72 (1968) 2782. b

In Table 98 are listed some thermodynamic constants for the formation of polyhalide ions in aqueous solution; with but few exceptions, however, the limitations of the physicochemical methods of analysis tend to impair the reliability of these data. For example, in calculations of the formation constant Ä> of a trihalide ion, it has commonly been assumed that the activity coefficients of the mono- and trihalide ions are equal and therefore cancel out in the expression for Ä>; in fact, this simplifying assumption is not valid. Still more fallible are the thermodynamic properties deduced for polyhalide ions containing more than three atoms, since the existence of successive equilibria is typically characterized by comparatively small departures from the pattern expected of the initial process XY+Z~v^XYZ~. Hence, the stabilities of most pentahalide species are still poorly defined, while the existence of other species remains problematic; falling within the second

1541

POLYHALIDE ANIONS 2

BrCl65

-, IC165 category, for example, are the ions I 6 - and the somewhat implausible 3_ and IBr 4 , which have been variously postulated as ingredients of aqueous equilibria840. Studies of some of the mixed polyhalide ions are further complicated by their tendency to disproportionate in favour of symmetrical species, e.g. 2BrICl" -> IBr2" +IC1 2 - 84 ° 2Br 2 Cl- -> Br 3 " + BrCl 2 - 975

As with solid polyhalide salts, dissociation in solution invariably takes the course which yields the smallest monohalide ion. Moreover, the stabilities of the polyhalide ions also follow a trend which is similar in outline to that established for the solid derivatives, though, not surprisingly, the intervention of a solvent appears to produce some differences of detail between the two stability sequences. The transition from water to a non-aqueous solvent is commonly attended by a dramatic increase in the stability constant of the poly­ halide ion. Set against a value of 710 in aqueous solution, for example, the formation constant of the Ϊ3~ ion at 25°C increases by several orders of magnitude when water is replaced as the solvent by 1,2-dichloroethane (Kf = 107), nitromethane (Kf = 106'7), acetone (Kf = 108*3) or acetonitrile (Kf = 106'6); likewise Kf for the Br3~ ion is reported to take the following values in different solvents at 25°C: 16 (water), 177 (methanol) and 1-66 xlO 4 (bromotrichloromethane). The enhanced stability of polyhalide ions in non-aqueous media presumably accounts also for the claim laid to the existence of the Cl5 anion in acetonitrile solutions for which the ratio [Cl2]/[Q-] exceeds unity976; as yet, this ion is unknown in aqueous solution. There is evidence, too, that the relative stabilities of different polyhalide ions varies from solvent to solvent: for example, when nitromethane, acetone or acetonitrile is the solvent, the trihalide ions appear to gain in stability in the sequence 13" < B ^ - < Cl3~, which is the reverse of that found in aqueous solution. Similarly, the stability of the IC1 4 - ion in solution depends on the nature of the solvent: in acetonitrile or 1,2-dichloroethane it dissociates to ICI2" and CI2 (Kf = ca. 7000 or 5400, respectively), whereas in trichloroacetic acid it dissociates completely to IC1, CI2 and Cl" 84°. Variations of stability of a particular anion have been correlated with the dielectric constant of the medium, although, since the formation of a polyhalide ion is not accompanied by the separation of charges, there is no prima facie reason for such a correlation. It seems more likely that the variations are attributable, at least in the first instance, to the protean effects of solvation and to the competition introduced by the acidic or basic functions of the solvent molecules. Discussion^11 The mode of decomposition of solid polyhalide salts and the dependence of the thermal stability on the size of the cation can be interpreted by thermodynamic reasoning analogous to that presented, in general, for other thermally unstable complex halides (Section 3, p. 1251) and, in particular, for solid derivatives of the hydrogen dihalide anions (Section 3, p. 1319). Thus, the factors which influence the mode of decomposition of a mixed polyhalide M+XYZ- may be established by considering the alternative decomposition paths and constructing the cycle depicted in Fig. 37. 975 R. p. Bell and M. Pring, / . Chem. Soc. (A) (1966) 1607. 976 j . c . Evans and G. Y.-S. Lo, / . Chem. Phys. 44 (1966) 3638. 977 A. G. Sharpe, Halogen Chemistry (ed. V. Gutmann), Vol. 1, p. 29. Academic Press (1967); D. A. Johnson, Some Thermodynamic Aspects of Inorganic Chemistry, p. 59, Cambridge (1968).

1542 CHLORINE, BROMINE, IODINE AND ASTATINE: A. J. DOWNS AND C. J. ADAMS ΔΗ

MXYZ (s)

MWUY7M (s) (g)

V^M^) + X - ( g ) ^ ^ M ^ ( g ) + X(g) ■" +YZ(g) ""+ γ ζ ω + c N s &

>MY(?)+X7(e)

__U* + 2RT M+(g) + Y"(g) V 5'2RT M+(g) + Y(g) D2 "-"+XZ(g) - + XZ(g) + e '

ι ><

AH2

k

MZ(s)+XY(g).

V

2RT

M+(g) +X(g) + Y(g)+Z(g)

M+(g) +Z-(g) E 3 + 5/2RT M*(g) + Z(g)^< + XY(g) - +XY(g) + e '

FIG. 37. Modes of decomposition of a mixed polyhalide M + XYZ~. That the overall change in entropy varies but little for the different paths is evident from the fact that a reaction such as KI(s) + £Cl 2 (g) -> KCl(s) + £I 2 (g) experiences an entropy change of less than 1 cal d e g - 1 . Variations of free energy are therefore determined almost exclu­ sively by variations in the enthalpy terms Δ / f j , Δ / / 2 and AH3. Reference to the cycle then shows that it is the interplay of the differences (a) between the lattice energies of the monohalides (Ui9U2 or C/3), (b) between the electron affinities of the halogen atoms (Eh E2 or £3), and (c) between the dissociation energies of the diatomic halogen molecules (D\9 D2 or D 3 ) that determines the course o f the reaction. In fact, the contributions o f (b) and (c) are mutually compensating, so that, overall, the determining factor in the decomposition is the lattice energy of the solid monohalide, which increases as the anion becomes smaller. Thus, for the complex CsIBrCl, the situation is as follows: (i) CsI(s) + BrCl(g): Di + Ex = 124 kcal and Ux = 144 kcal; (ii) CsBr(s)+ICl(g): D2+E2 = 129 kcal and U2 = 151 kcal; and (iii) CsCl(s)+IBr(g): 2 > 3 + £ 3 = 127 kcal and £/ 3 = 156 kcal. Hence, the sequence ΔΤ/i > ΔΗ2 > A / / 3 is implied, the favoured products being CsCl(s)+IBr(g). Even if the decomposition products include the diatomic halogen in the condensed phase, the additional enthalpy terms are still transcended by the differences in the lattice energies of the monohalides; moreover, the effects of condensation are further mitigated by an unfavourable entropy change. Thus, in contrast with solid derivatives o f the hydrogen dihalide anions, which decompose to give the hydrogen halide with the maximum bond energy, solid polyhalides decompose to give the monohalide having the smallest anion. The difference in behaviour between complexes of the two classes reflects the circumstance that changes in the nature of the halogen evoke substantial variations in the bond energies of the hydrogen halides but only modest variations in those of diatomic halogen or interhalogen molecules. A n interpretation of the mode of dissociation of polyhalides in solution follows similar lines, except that the production of the smallest anion is then dictated by the solvation energies rather than the lattice energies of the different species. The effect o f cation size o n the stability of solid derivatives of a particular polyhalide anion, e.g. IC1 2 ~, may be discussed in terms of the following thermodynamic cycle: M+[IC1 ]-(·)

£i£

-UlMCl]-2RT

U[MIC1]+2RT M+(g)

-M+Cl"(s) + I d (g)

(ια2Γ(8)-

—

-M + (g) + Cf"(g) + ICl(g)

If variations in the nature o f the cation M + are assumed, t o a first approximation, not t o

1543 affect the entropy term associated with the decomposition reaction, the pattern of thermodynamic stability is inevitably determined by the enthalpy change ΔΗ°. According to the cycle, POLYHALIDE ANIONS

AH° = £/[MICl2]-tf[MCl]+*

where x is a constant for a given polyhalide anion. Using Kapustinskii's approximation to express the lattice energies of the simple and complex halides in terms of the appropriate ionic radii r, we have (in kcal) Lr(M+)+r(ICl2-)

r(M+)+r(C\-)l

Since the effective thermochemical radius of the [IC12] ~ anion must exceed that of the Cl anion, it follows that the contribution to ΔΗ° made by the lattice energies is intrinsically exothermic. However, the incentive to decomposition represented by this term clearly decreases as the radius of the cation expands. The ionic model therefore accounts success­ fully for the observation that solid derivatives of a particular polyhalide anion become pro­ gressively more stable as the size of the cation increases. The limitations of multivalent cations Mn + in stabilizing polyhalide anions depend, not only on the smaller size of such cations, but also on the heightened differences between the lattice energies of the simple and complex halides, since for the decomposition of 1 g-ion of ICI2 ~ in the salt M(ICl2)w ™° =

256(

" + 1 ) [ K M » + ) U C I 2 - ) - KM»+HKCI-)] + *

Another contributory factor is probably the destabilizing effect of anion-anion contact, the impact of which is most severe on derivatives of multivalent cations; in fact, it is likely that the anions in complexes like Ba(BrF4)2 take the form of extended polymeric networks. Although the effectiveness of large tetra-alkylammonium ions like NMe 4 + in forming stable polyhalides appears to comply with these principles, the recently measured free energies of decomposition of some polyhalide salts [NR4]I» show that the corresponding entropy changes are both opposite in sign and larger in magnitude than those observed for simpler cations like Rb +, Cs + and NH4 + 974. According to the results of Table 97, the disordering which accompanies the formation of a solid tetra-alkylammonium polyiodide makes a major contribution to the stability of the complex. On the basis of these and other thermodynamic parameters, the following enthalpy terms (at 25°C) have been deduced: ΔΗ° = -24-0kcal for the reaction l2(g)+I~(g)-> h~(g)l Atf/tVfe)] = -56- 4 kcal; and Δ//° = —44·0 kcal for the hydration of the gaseous I3 - ion974. Although strictly a net en­ thalpy change, the first of these terms is analogous to the so-called "hydrogen bond energy" of an HX2 ~ anion977, testifying indirectly to the strength of the bonds in the iso­ lated tri-iodide anion. 2. Structures of solid polyhalides834»835»840»978. Our present knowledge of the struc­ tures of polyhalide anions has been gained pre-eminently from X-ray diffraction studies of crystalline salts, though these have lately been augmented by recourse to neutron diffraction and vibrational, nqr and Mössbauer spectroscopy. Summarized in Table 99 are the struc­ tural details presently available for anionic polyhalide derivatives; the geometries of the anions are also depicted in Fig. 38. 978 R. E. Rundle, Ada Cryst. 14 (1961) 585.

1. Trihalide species

(a)

0-0—O

0»

o—o-o Unsymmetrical

Symmetrical

2. Pentahalide species ^2-82 Ä

3-17 X

0>)

(a)

I5" in Me4NI5

BrF4"*in KBrF4and ia;inKICl 4 ,H 2 0

3. Heptahalide species 80-3^

80-3°

a=2.904 Ä b=3.435 Ä c=2.735 A

Structure of anionic network in Et4NI7: I'ions and I2 molecules are linked to form a three-dimensional array

4. Enneahalide species 168°

5. Other species 116°

178" 3-24 A ,

o^^PQao 0

3-18Ä

2-90Ä

300 X

2-85 A

Structure of anionic network in Μβ4Νΐ9 Structure of the I 8 2 ~ anion in CS2I8 FIG. 38. Structures of anionic polyhalide aggregates.

1544

I3-

Br3-

Cl3-

Modification I: contains two indepen­ dent centrosymmetric I3~ ions (X)15 Modification II: contains two inde­ pendent nearly linear but unsym­ metrical I3" ions (X)!5

The structure contains I3~ ions which appear to be linear and symmetrical, alternating with I~ ions along the trigonal axis of the unit cell (X)16

NEt4+

+

K in KI,KI3,6MeCONHMe

Nearly linear, symmetrical (X) Nearly linear, slightly unsymmetrical I3~ ions linked in chains (X)14

AsPlu [PhCONH2]2H+ in 2PhCONH2,HI3

7

Nearly linear, unsymmetrical (X)7

NH 4 +

+

Nearly linear, unsymmetrical (X)7

Nearly linear, symmetrical (X)

7

Me3NH in [Me3NH+]2Br-Br3Cs+

+

Nearly linear, unsymmetrical (X)6

PBr4+

Linear, symmetrical (V)

Shape of anion

Nearly linear, unsymmetrical (X)2

Cation

NEt 4 \NPr n 4 + orNBun4+ Cs+

TRIATOMIC ANIONS

anion

Polyhalide

= = = = =

2-82 310 2-90 2-910 2-951

(a) I-I = 2-928 (b) I-I = 2-943 (a) I-I = 2-912, 2-961 (b) I-I = 2-892, 2-981 I-I = 2-945

I-I I-I I-I I-I I-I

Br-Br = 2-440 Br-Br = 2-698 Br-Br = 2-39 Br-Br = 2-91 Br-Br = 2-53, 2-54 I-I = 2-83 I-I = 3 03

(A)

Bond lengths

SOLID PHASE

ZI-I-I = 180°

ZI-I-I = 177-7° 2-666

ZI-I-I = 180°

Z.l-l-1 = 176° ZI-I-I = 176-8°

ZI-I-I = 177°

Z.I-I-I = 176°

Z.Br-Br-Br = 171°

ZBr-Br-Br = 177-3°

ZI-I-I = 179-5°

2-666

Vi

Spectroscopic studies

νβ-π NQR512 Mi3

ZBr-Br-Br = 177-5° V34NQR5

—

Bond angles

2-666 2-666 2-666

2-666 2-666 2-666 2-666 2-666

2-666 2-666

2-281 2-281 2-281 2-281 2-281

—

Bond length in correspond­ ing diatomic molecule (A)

TABLE 99. STRUCTURAL DETAILS FOR ANIONIC POLYHALIDE COMPLEXES

Cation

Cs +

NH 4

I2Br"

IBrCl"

+

Cl 5 Br 5 -

PENTA-ATON« c ANIONS

Cs

+

— —

piperazinium in [C4H12N2][IC12]2 triethylenediammonium in [HN(C2H4)3NH][IC12]2

PC14+

Rb + and Cs + Cs+,NMe4+, NEt4+,NPrV and N B u V NMe 4 +

NO+

IBr2"

ici2-

BrCl2~

CtfV

anion

Polyhalide

TABLE 99 (cont.)

— —

Nearly linear [Cl-I-Br]" ions statistically distributed in the crystal (studied at 140°K, X)24

Nearly linear, unsymmetrical [Br-I-Br]- ion (X)22 Nearly linear [I-I-Br]" ion (X)23

Structure contains two non-equivalent [Cl-I-Cl]" ions, both linear and unsymmetrical (X) 21

i o n (X)20

Linear, unsymmetrical [Cl-I-Cl]"

Linear, symmetrical [Cl-I-Cl]" ion (X) 19 Linear, symmetrical [Cl-I-Cl]" ion (X) 7

(V)3,18

Linear, symmetrical [F-Cl-F]" ion (V)n Unsymmetrical [F-Cl-F]" ion (V)*? Linear, symmetrical [Cl-Br-Cl]~ ion

Shape of anion

(A)

— —

I-Cl = 2-36 though probably in error I-Cl = 2-47 I-Cl = 2-69 (a) I-Cl = 2-54, 2-67 (b) I-Cl = 2-53, 2-63 I-Br = 2-62 I-Br = 2-78 I-I = 2-777 I-Br = 2-906 I-C1= ~2·91 I-Br = ~2-51

I-Cl = 2-55

—

—

Bond lengths

SOLID PHASE

— —

= 178-0°

— —

ZBr-I-Cl = 179°

ZI-I-Br

L Br-I-Br = 178-0°

L Cl-I-Cl = 180°

2-321 2-485 2-485 2-666 2-485 2-321 2-485

ZC1-I-C1 = 180°

L Cl-I-Cl = 180°

Z Cl-I-Cl = 180°

L Cl-I-Cl = 180°

—

—

Bond angles

2-321 2-321 2-321

2-321

2-321

—

—

Bond length in correspond­ ing diatomic molecule (A)

Vi V4

V8

V8.9 NQR12 V8.10

NQR5

V3.8

V3.18

V17

Spectroscopic studies

Cs +

K + in KIC1 4 ,H 2 0

IF4-

ICI4-

K+,Cs+,NMe4+ and NEt 4 +

IF 6 "

NMe 4 +

Ϊ82"

Cs +

OTHER SPECIES

V

Planar Z-shaped Is 2 " anions present, having two arms consisting of asymmetric 13" anions linked by an I2 molecule (X) 7

The structure can be considered to arise from the association of I7 " units with additional I2 molecules (X)7 (see Fig. 38)

Dia symmetry implied for the BrFe" anion (V) 32 ; rhombohedral crystal system (X) Non-octahedral symmetry implied for the IF 6 " anion (V and M)33,34

K + , R b + orCs+

BrF 6 "

ENNEA-ATOMIC ANIONS

Structure contains centrosymmetric I3" anions linked to I2 molecules in a 3-dimensional network (X)? (see Fig. 38)

NEt 4 +

I7-

HEPTA-ATOMIC ANIONS

Approximately square-planar IC14~ anion (X)30

K+

BrF4"

(V) 2 9

Civ symmetry implied for the anion

Rb + and Cs +

C1F4-

Nearly planar V-shaped I5" anions (X)7 (see Fig. 38) Square-planar C1F4" anion with D*h symmetry (V) 25 Square-planar BrF4~ anion with Ζ)4Λ symmetry (N) 2 6

NMe 4 +

ι5-

I-I = 2-85 and 3 0 0 (I3-) I-I = 2-80 (I 2 ) and 3-42 for shortest I2· Ί 3 " distance

I-I = 2-67, 2-90, 3-18, 3-24 and 3-43

I-I = 2-904 (I 3 -), 2-735 (I 2 ) and 3-435 for shortest I2· I3" distance

I-Cl = 2-42, 2-47, 2-53 and 2-60

—

Br-F = 1-89

—

Iapex-I = 3-17 Iterminal-1 = 2*82

Z I - I - I = 177° (I 3 -) Ζ Ι - Ι - Ι - · Ι = 81° Z.I 1-1= 175°

2-666

zlI-I-I = 85,87 and 89°; 156, 168, 169, 175 and 178°


2-666

2-666

2-666

ZCl-I-Cl(adj) = 90-6, 90-7, 89-2 and 89-5°

—

— 2-321

/.F-Br-F(adj) = 90-9, 891 ±2-0°

ZLI-Iapex-I = 95° Z.Iap-1-Iterm = 175°

1-756

—

2-666 2-666

NQRi 2

yio

V33 M34

V32

yio

V3 NQR5 M3i

Λ/29

V27-29

V25

yio

M = Mössbauer spectra; N = neutron diffraction; NQR = nqr spectra; V = vibrational spectra; X = X-ray diffraction, i J. C. Evans and G. Y.-S. Lo, / . Chem. Phys. 44 (1966) 3638. 2 G. L. Breneman and R. D. Willett, Acta Cryst. B25 (1969) 1073. 3 W. B. Person, G. R. Anderson, J. N . Fordemwalt, H. Stammreich and R. Forneris, / . Chem. Phys. 35 (1961) 908. 4 J. C. Evans and G. Y.-S. Lo, Inorg. Chem. 6 (1967) 1483. 5 E. A. C. Lücken, Nuclear Quadrupole Coupling Constants, p. 289. Academic Press (1969). 6 G. L. Breneman and R. D . Willett, Acta Cryst. 23 (1967) 467. 7 E. H. Wiebenga, E. E. Havinga and K. H. Boswijk, Adv. Inorg. Chem. Radiochem. 3 (1961) 148. 8 A. G. Maki and R. Forneris, Spectrochim. Acta, 23A (1967) 867. 9 G. C. Hayward and P. J. Hendra, Spectrochim. Acta, 23A (1967) 2309. io F . W. Parrett and N. J. Taylor, / . Inorg. Nuclear Chem. 32 (1970) 2458. ii S. G. W. Ginn and J. L. Wood, Chem. Comm. (1965) 262. 12 G. A. Bowmaker and S. Hacobian, Austral. J. Chem. 21 (1968) 551. 13 B. S. Ehrlich and M. Kaplan, / . Chem. Phys. 51 (1969) 603. 14 J. M. Reddy, K. Knox and M. B. Robin, / . Chem. Phys. 40 (1964) 1082. is T. Migchelsen and A. Vos, Acta Cryst. 23 (1967) 796. 16 K. Toman, J. Honzl and J. Jecny, Acta Cryst. 18 (1965) 673. 17 K. O. Christe, W. Sawodny and J. P. Guertin, Inorg. Chem. 6 (1967) 1159. is J. C. Evans and G. Y.-S. Lo, / . Chem. Phys. 44 (1966) 4356. 19 G. J. Visser and A. Vos, Acta Cryst. 17 (1964) 1336. 20 Chr. Romming, Acta Chem. Scand. 12 (1958) 668. 2i S. Gran and Chr. Romming, Acta Chem. Scand. 22 (1968) 1686. 22 J. E. Davies and E. K. Nunn, Chem. Comm. (1969) 1374. 23 G. B. Carpenter, Acta Cryst. 20 (1966) 330. 24 T. Migchelsen and A. Vos, Acta Cryst. 22 (1967) 812. 25 K. O. Christe and W. Sawodny, Z. anorg. Chem. 374 (1970) 306 and references cited therein. 26 A. J. Edwards and G. R. Jones, / . Chem. Soc. (A) (1969) 1936. 27 K. O. Christe and C. J. Schack, Inorg. Chem. 9 (1970) 1852 and references cited therein. 28 T. Surles, H. H. Hyman, L. A. Quarterman and A. I. Popov, Inorg. Chem. 9 (1970) 2726; ibid. 10 (1971) 913. 29 J. Shamir and I. Yaroslavsky, IsraelJ. Chem. 7 (1969) 495. 30 R. J. Elema, J. L. de Boer and A. Vos, Acta Cryst. 16 (1963) 243. 31 D. W. Hafemeister, The Mössbauer Effect and its Application in Chemistry, p. 126. Advances in Chemistry Series No. 68, American Chemical Society (1967). 32 R. Bougon, P. Charpin and J. Soriano, Compt. rend. 272C (1971) 565. 33 K. O. Christe, J. P. Guertin and W. Sawodny, Inorg. Chem. 7 (1968) 626; S. P. Beaton, D . W. A. Sharp, A. J. Perkins, I. Sheft, H. H. Hyman and K. Christe, ibid. p. 2174; H. Klamm, H. Meinert, P. Reich and K. Witke, Z. Chem. 8 (1968) 393, 469. 34 S. Bukshpan, J. Soriano and J. Shamir, Chem. Phys. Letters, 4 (1969) 241.

TABLE 99 (cont.)

POLYHALIDE ANIONS

1549

On the basis of this information, polyhalide anions are evidently characterized by the following general features: (a) Heteronuclear anions, e.g. IC12 ~, IBrCl ~or IC14 ~, invariably adopt the form in which the halogen of highest atomic number is the central "polyvalent" atom. (b) Interatomic distances are always greater, by some tenths of an angstrom, than those in the corresponding diatomic halogen molecules. Apparent exceptions to this rule have now been diagnosed as the outcome of inaccuracies in earlier crystallographic studies of polyhalide salts containing the anions IC1 2 - 979 , ICl4-943 a n d BrICl-980. Though still substantially less than twice the van der Waals' radius of iodine, some of the I · · · I distances in the polyiodide salts NEt4I7, NMe^p and Cs2l8 are not much shorter than those between iodine molecules in crystalline iodine. (c) Bond angles close to 90° or 180° predominate in the anionic aggregates, though values as low as 80° or 156° have been encountered. (d) A trihalide ion varies its symmetry and dimensions according to the demands of its environment, the individual bond distances changing in a systematic way with the overall length of the ion978»981. This stereochemicalflexibilityis probably inherent, to some degree, in all polyhalide anions. Trihalides All of the trihalide anions which have been structurally characterized in the crystalline phase comply with the general rules (a)-(c): the anions are invariably linear, or nearly linear, with bond angles at the central atom in the range 171-180°. The characteristic (d) has been shown to apply to the anions Br3 _ , I3 - and IC12 ~, each of which has been studied in three or more salts. Trihalide ions may be either symmetrical or unsymmetrical in form. With the I2Br~ and BrICI" anions, found, respectively, in the salts CsI2Br982 and NH4BrICl980, the dissymmetry is a necessary consequence of the composition of the ion or of the operation of generalization (a). However, other anions, e.g. Br3 _ , I 3 " and IC12 ~, have the capacity to assume either a symmetrical or an unsymmetrical form, depending on the precise nature of the salt in which they occur. Thus, the Br3 - anion appears to be sym­ metrical in the mixed salt [Me 3 NH + ] 2 Br-[Br 3 ]~ 835 but unsymmetrical in CsBr3981 and [PBr4] +[Br3] - 973 ; the I3 - anion appears to be symmetrical in crystalline [AsPh4] + [I 3 ] ~ 835 , [NEt4]I7835, the complex KI,KI3,6MeCONHMe983 and modification I of [NEt4] +[I3] - 984, but unsymmetrical in Csl 3 8 ^, NH 4 I 3 835 , Cs 2 I 8 835 , the complex HI3,2benzamide985 and modification II of [NEt4] + [I 3 ] - 984 ; whereas the IC12 - anion is reported to be symmetrical in [NMe4]+[IC12] ~ 979 but unsymmetrical in the piperazinium986 and triethylenediammonium987 salts. The IBr2 - anion is unsymmetrical in the one salt whose structure has been determined, viz. CsIBr2988. In some cases, e.g. that of the complex KI,KI3,6MeCONHMe, it is impossible to rule out the circumstance that the appearance of centrosymmetry is determined by statistical disordering of the trihalide anions which are 979 G . J. Visser and A . Vos, Acta Cryst. 17 (1964) 1336. 980 T . Migchelsen and A . Vos, Acta Cryst. 2 2 (1967) 812. 981 G . L. Breneman and R. D . Willett, Acta Cryst. B25 (1969) 1073. 982 G . B . Carpenter, Acta Cryst. 20 (1966) 330. 983 K . T o m a n , J . Honzl a n d J. Jecny, Acta Cryst. 18 (1965) 673. 984 T . Migchelsen and A . Vos, Acta Cryst. 2 3 (1967) 796 985 J . M . Reddy, K . K n o x a n d M . B . Robin, / . Chem. Phys. 40 (1964) 1082. 986 Chr. Romming, Acta Chem. Scand. 12 (1958) 668. 987 s . G r a n and Chr. Romming, Acta Chem. Scand. 2 2 (1968) 1686. 988 j . E . Davies and E . K . N u n n , Chem. Comm. (1969) 1374.

1550

CHLORINE, BROMINE, IODINE AND ASTATINE: A. J. DOWNS AND C. J. ADAMS

individually unsymmetrical. The general pattern of behaviour is that centrosymmetric trihalide anions are favoured by the presence of bulky cations such asNR4+or AsPh4 + , though the recognition of two structural modifications of the salt [NEt4][l3]984, one containing symmetrical and the other unsymmetrical I3 ~ anions, implies that the size of the cation is not the only factor critical to the shape of the anion. In general, the variations of structure are attributable to the electrostatic field developed by ions in the neighbourhood of a given trihalide anion978»984»989. Molecular-orbital calculations which take account of this crystal field for the I3 - ion989 imply that one equilibrium position is available to the central iodine atom when the distance between the two terminal iodine atoms D becomes equal to or less than a certain critical value D c , whereas there are two such positions for this atom when D exceeds Dc (see Fig. 39). When D is large, the I 2 * · * I" interaction is essentially that of an 1

1

11

1

11

««*

«***

«" "* **

110

"^

^

*" " ""

^ -- ""

yf

Λ^


■-' -**

^*

Χ^

--

^^ ^^

*"^ -"χ ^^y^^»^

cL limit I ~ 2

1\ J

\*0*^\

^ ^

^ * ^ * ^^

I3

0

^^ "^ ^^

> ^

^

■?

^ ^ Β Γ Γ <**^ 3

**

y

^

^" tyimit Br3~

Line of centres Br3~

_ ^ - = : ^

j

Line of centres I ~ 3

5*»»—"V" '

"j

^

ΙΓ Br

—"—■

r

d, limit 1 3 dj limit Br3~ 0-90

;

i

i

i

1000

1020

1040

i

D/D c

FIG. 39. Comparison of the configurations of the Br3 ~ and I3" anions (after ref. 981). D = total length of the X3" ion; Dc = critical value of D where the ion becomes sym­ metrical; dv = Xoentrai-Xterminai distance with i = 1 for the short bond and i = 2 for the long bond.

ion with a polarizable molecule, but as D approaches Z)c, orbital-overlap and covalent bond­ ing become increasingly important. These general features disclose a striking resemblance between the triatomic systems X2-X" and H2-H978»981, though with the major difference 989 R . D . Brown and E. K. Nunn, Austral J. Chem. 19 (1966) 1567.

POLYHALIDE ANIONS

1551

that the magnitude of D for the former is acutely responsive to the influence of neighbour­ ing ions. In a salt containing relatively small counter-ions, e.g. Cs + , the crystal field experi­ enced by the I3 ~ ions is unsymmetrical, with the result that one of the two equilibrium posi­ tions open to the central iodine atom is preferentially stabilized; consequently the I3 "* assumes an unsymmetrical form. By contrast, the relatively weak, symmetrical crystal field created by an environment of large cations, e.g. AsPh4+, produces minimal perturbation of the short, centrosymmetric unit which characterizes the isolated I3 - ion. Somewhat less plausible is the suggestion of Slater990 that D for the isolated I3 ~ ion is comparable with the largest values observed in crystals, but that the "pressure" of certain environments forces the ion to become shorter and, ultimately, centrosymmetric. Unfortunately, no clearcut decision is possible concerning the shape of trihalide ions in solution. Whereas the infrared spectrum of the salt [NBu4][I3] in benzene, nitrobenzene and pyridine is consistent with the presence of centrosymmetric I3 ~ ions991, the Raman spectra of solutions containing Br3 _ ,I 3 ~ or IBr2 -anions do not comply with such an assump­ tion992. In view of these ambiguities, it is impossible, as yet, to assess the effect of the anisotropic environment which the ions are likely to experience in solution. The kinship of the trihalide ions with charge-transfer complexes formed between diatomic halogen or interhalogen molecules and neutral donor species is underlined by reference to the "effective radii" R\ and R2 of the central iodine atom in units of the type X-I-Y (X = halogen; Y = halide ion or donor atom of neutral molecule), these being evaluated by subtracting the accepted covalent radii of the ligand atoms X and Y from the measured interatomic distances993. To a good approximation, a linear relationship is found to connect R\ and R2, viz. R2 = -1-97/?!+4-59 Ä, irrespective of whether X-I-Y corresponds to a trihalide anion or to a neutral complex. Quantitative expression is thus given to the generalization that the strengths of the bonds formed by the central atom are complementary, the intermolecular I-Y bond gaining in strength at the expense of the intra­ molecular I-X bond until, as with the centrosymmetric I3 ~ anion, the distinction between the two bonds is lost. Higher Polyhalides The structures of pentahalide ions and of more complicated aggregates do not show the same regularity as those of the trihalide anions. The three pentahalide anions for which definitive results are available adopt one of two structures: (i) a square-planar array, or (ii) a nearly planar V-shaped unit (see Fig. 38). Following earlier vicissitudes, the squareplanar structure of the BrF4~ anion in the salt KBrF4 has now been established by neutron diffraction937, while refinement of the crystal structure of the salt KICl4,H20943 reveals a similar, though somewhat distorted, structure for the IC14 - ion. By contrast, the I 5 - ion in the complex [NMe4][I5] belongs to the second category: each arm of the nearly planar V-shaped unit corresponds to an unsymmetrical 13" anion, the apex angle being 95°; the ion may thus be formulated as [1(12)2] ~ with two iodine molecules coordinated to a single iodide ion. On the premises of the vibrational spectra, the B r F 4 - 9 3 4 - 9 3 6 and ICI4- 994 ions retain their square-planar geometry in solution, and are isostructural with the 990 j . c . Slater, Ada Cryst. 12 (1959) 197. 991 S. G. W. Ginn and J. L. W o o d , Chem. Comm. (1965) 262. 992 A . G. Maki and R. Forneris, Spectrochim. Ada, 23A (1967) 867. 993 o . Hassel and Chr. Romming, Ada Chem. Scand. 21 (1967) 2659. 994 w . B. Person, G. R. Anderson, J. N . Fordemwalt, H . Stammreich and R. Forneris, / . Chem. 35 (1961) 908.

C.I.C. VOL I I - B B B

Phys.

1552

CHLORINE, BROMINE, IODINE AND ASTATINE: A. J. DOWNS AND C. J. ADAMS

C1F4" ion in its rubidium or caesium salt932. On the other hand, a V-shaped skeleton analogous to that of Is" may well be assumed by the species Cl 5 " 976 , Br 5 " 995,12CI3" 996 and I2Cl2Br"996, a view consistent with the incompletely defined vibrational spectra attrib­ uted to the anions, but as yet unsubstantiated by more compelling evidence. The complex [NEt4]I7 is the only representative of the group of heptahalides which has been the subject of detailed structural analysis835»840. The crystal of this compound contains, not discrete I 7 - ions, but a three-dimensional network composed of centrosymmetric 13" ions and I 2 molecules; adjacent to each I 3 - ion are four I 2 molecules, and to each I 2 molecule are two 13 - ions (Fig. 38). Hence, the complex is more realistically formulated as [NEt 4 ] + [I 3 (I 2 ) 2 ]-. At 3-435 A the closest approach of the I 3 ~ and I 2 units differs but little from the shortest intermolecular distance (3-50 A) in crystalline iodine; that it is substantially less than twice the van der Waals' radius of iodine (4-30 A) indicates, nonetheless, comparatively strong interaction between the I 3 - and I 2 units. The structures adopted by the heptahalide anions BrF 6 - and IF 6 _ are of especial interest in view of their relationship to xenon hexafluoride997, with the implication of a valence shell for the central atom having a complement of 14 electrons. Vibrational946b»949 and 129I Mössbauer950 spectra signify non-octahedral structures for both anions whether in solution or in the crystalline phase; it has not been possible to specify the precise geometry, though the vibrational spectra of the crystalline complexes MBrF 6 (M = K, Rb or Cs) imply that the octahedron of fluorine atoms is distorted about a threefold axis to impart D$d symmetry to the environment of the bromine atom 946b . The crystal structure of the salt [NMe4]I9 reveals aggregates of iodine atoms which may be considered as I 9 - anions, though one of the interatomic distances (3-43 A) is scarcely shorter than the closest approach between iodine atoms of different I 9 ~ units (3-49 A)835»840. It is perhaps more realistic to consider the structure in terms of non-planar Z-shaped I 7 - ions each linked by relatively strong interaction with an additional I 2 molecule and with other I 7 - entities. A more distinct anionic unit is observed in the crystal structure of the salt formulated on the strength of its observed diamagnetism as Cs2I8835»840. Here it is possible to distinguish planar Z-shaped I 8 2 - aggregates composed of two unsymmetrical I 3 - ions linked by a common I 2 molecule. To be represented therefore as [I2(l3)2]2~> each aggregate involves interaction between the I 3 - ions and I 2 molecule (I · · · I distance 3-42 A) which is substantially stronger than that between adjacent I 8 2 _ aggregates (shortest I · · · I approach 3-88 A). 3. Spectroscopic properties. The thermodynamic and structural investigations outlined in the preceding sections have been augmented by measurements relating to the vibrational992»994, electronic84***998 -ιοοο, nqr9i7,iooi a nd i27I or *29I Mössbauer95
1553

POLYHALIDE ANIONS

Thus, the infrared and Raman spectra of polyhalides have been employed, not only to investigate the structure of the anions in different environments, but also to determine the properties of the potential field experienced by the atoms within the anion. Valence force constants calculated from the vibrational frequencies of polyhalide anions known or presumed to be symmetrical are listed in Table 100; particular interest attaches to the TABLE 100. FORCE CONSTANTS FOR SYMMETRICAL POLYHALIDE ANIONS

Anion

ci 3 -

a

Br 3 -* l3"C

ClF 2 - d BrCV e IQT*·'

IBr 2 - f ·* ClF 4 - h BrF4"h IF 4 ~ h

ici 4 - b

Principal stretching force constant, Λ(ΥΧΥ) (mdyne/A)

Interaction constant between collinear

0-96 0-91 0-71 2-35 102 100 0-91 2-13 2-23 2-22 1-25

0-55 0-32 0-23 017 0-47 0-36 0-30 0-23 0-20 018 0-33

bonds, /rr(YXY)

(mdyne/A)

Interaction Bending force constant constant for between perpendicular linear YXY bonds, /rr'(YXY)

unit,

0-57 0-36 0-32 007 0-46 0-36 0-33 011 009 008 0-26

0-30 0-38 0-42 0-55 0-36 0-46 0-42 0·475 0·545 0-62 0-53

Air*

(mdyne/A)

(mdyne/A)

— — — — — — —

— — — —

0-27 0-43 0-47 008

/rr(YXY) /r(YXY)

Λ(ΥΧΥ) MXY) L/KXY) = stretching force constant of XY molecule]

013 010

— — — —

a

J. C. Evans and G. Y.-S. Lo, / . Chem. Phys. 44 (1966) 3638. W. B. Person, G. R. Anderson, J. N. Fordemwalt, H. Stammreich and R. Forneris, / . Chem. Phys. 35 (1961)908. c S. G. W. Ginn and J. L. Wood, Chem. Comm. (1965) 262. d K. O. Christe, W. Sawodny and J. P. Guertin, lnorg. Chem. 6 (1967) 1159. e J. C. Evans and G. Y.-S. Lo, J. Chem. Phys. 44 (1966) 4356. f A. G. Maki and R. Forneris, Spectrochim. Acta, 23A (1967) 867. β G. C. Hayward and P. J. Hendra, Spectrochim. Ada, 23A (1967) 2309. h K. O. Christe and D. Naumann, lnorg. Chem. 12 (1973) 59. b

comparison of the principal stretching force constant fr(YXY) of each anion with the corresponding interaction constant between collinear bonds^.r(YXY).and with the stretching force constant of the corresponding diatomic halogen molecule fr(XY)· Analysis reveals, firstly, that the force constant/r(YXY) is only 30-58% of the magnitude of/ r (XY), and, secondly, that the interaction constant frr(YXY) commonly takes unusually large values in proportion to^(YXY); as the ratio frr(YXY)/fr(YXY) increases over the range 0Ό7-0-57, so fr(YXY)/fr(XY) diminishes. These findings confirm the analogy between the polyhalide and hydrogen dihalide anions, both being characterized by unusually weak bonds; further, in that the term [fr(YXY)-frr(YXY)] defines the force constant for vibration of the central atom within the one-dimensional potential well bounded by the ligands, the curvature of the well is singularly flat at the equilibrium position of the central atom. That the central atom of such an anion is thus relatively loosely confined to its equilibrium position accounts for the susceptibility to structural changes of the anion under the influence of different crystal or solvation forces. The marked inferiority of the valence force constants of the polyhalide anions, compared with corresponding diatomic halogen molecules or with polyhalogen

1554

CHLORINE, BROMINE, IODINE AND ASTATINE: A. J. DOWNS AND C. J. ADAMS

cations such as C1F2 + , supports the molecular-orbital scheme (see Fig. 4) first developed by Pimentel and Hach and Rundle1003, in which only the valence ^-orbitals of the halogen atoms are presumed to subscribe to the bonding. The results may then be reconciled with the three-centre, four-electron character of the bonds in polyhalide anions; approximate bond orders of 0-5 or less may be divined via the ratio fr(YXY)/fr(XY). Variously exploited for the measurement of stability constants, the optical spectra of polyhalide anions in aqueous or non-aqueous media have been described by several investigators840»998 -ιοοο^ but with the interpretation of the details little headway has been made. With the assumption of valence-shell molecular orbitals composed exclusively of atomic /7-functions, the molecular orbital scheme for a centrosymmetric trihalide anion takes the form depicted in Fig. 40, whence the two transitions of lowest energy are seen to be

^

Anti-bonding i

Forbidden

Allowed

L_ H ■ | |

j!

=

\ Non-bonding

f

Bonding

■ # ■

FIG. 40. Molecular-orbital scheme and electronic transitions for a linear trihalide anion.

au* <- TTU* and ση* «- ag. As with the parent halogen molecules, the first of these is formally forbidden in a centrosymmetric unit; that it appears relatively weakly in absorption at 23,000-30,000 cm - 1 1 0 0 ° may be due partly to spin-orbit coupling phenomena and partly to the distortion impressed upon the ion by crystal or solvent interactions. The transition occurs, in effect, within the /?-shell of the central atom, and the associated absorp­ tion has some of the characteristics of a ligand-field band; the following frequencies (in cm _1 ) show signs of the evolution of a spectrochemical series: I 3 - , 23,000; IBr 2 ~, 27,000; IBrCl - , 1003 G . C. Pimentel, / . Chem. Phys. 19 (1951) 446; R. J. Hach and R. E. Rundle, / . Amer, Chem. Soc. 73(1951)4321.

1555

POLYHALIDE ANIONS

28,100; IC1 2 -, 29,800. On the other hand, the transition au**-ag, which is allowed irrespective of the geometry of the trihalide anion, entails electron-transfer from one of the terminal atoms to the central halogen atom: [ Y — x — Y ] - - V Ä ~ . [γ—χ-

γ]

a process plainly akin to the charge-transfer which characterizes neutral complexes of the diatomic halogen molecules. Analogies with the spectra of such complexes, combined with thermodynamic arguments1000, favour the assignment to this transition of the intense absorption bands observed at 28,000-44,000 cm - 1 . Spin-orbit splitting is then presumed to account for the intense doublet feature (at 28,300 and 34,800 cm - 1 ), to which the aqueous 13 ~ anion owes its red colour. Alternatively the bands may have their origin in the transi­ tions au* <- TTU* and ση* <-7Tg, having probabilities much enhanced by the mixing of σ- and ττ-orbitals which would accompany the departure from linearity of the I 3 ~ skeleton" 8 . If the latter interpretation is extended to a one-dimensional chain of I 3 ~ ions, subject to exciton-like coupling, a semi-quantitative explanation is found for the blue colour of the starch-iodine and related complexes998. The quadrupole coupling constants derived from the nqr917»1001 or iodine Mössbauer950»1002 spectra have been used to explore the electron distribution in the valence orbitals of the atoms forming polyhalide species. That the total ρ-σ electron population of a trihalide anion is invariably close to 4, within the limitations of the Townes-Dailey approximation, vindicates the four-electron, three-centre scheme of bonding; on similar grounds, a bonding scheme of this sort is also favoured by the quadrupole coupling constants for the pentahalide anion ICI4 ~. The essential /^-character of the bonding is endorsed, moreover, by the measured isomer shifts of the few Mössbauer spectra950»1002 which have been reported for iodine-containing polyhalide salts. 0-Decay of the 129I nucleus in the salts KIC1 4 ,H 2 0 and KIC1 2 ,H 2 0 has provided an ingenious method of synthesizing in situ the previously unknown compounds XeCl4 and XeCl2, which are short-lived under conventional conditions; since the decay populates the excited nuclear level of 129Xe, the products are conveniently detected and characterized by their Mössbauer spectra1004. The properties deduced for the molecules XeCU and XeCl2 stress their close resemblance, in bonding and structure, to the respective isoelectronic species IGU - and IC1 2 ~, from which they issue. Chemical Properties 840

The chemical reactions of a polyhalide anion are largely those of the dissociation products, and therefore reflect the stability of the complex with respect to dissociation. The function of polyhalide derivatives as sources of the corresponding halogen or interhalogen compounds is illustrated by the fluorinating action of solid tetrafluorochlorate(III) and tetrafluorobromate(III) salts838»841; thus, although bromine trifluoride is commonly presumed to be the active principle, it is notable that the fluorination of a number of metal oxides proceeds to completion more quickly with potassium tetrafluorobromate(III) than with the parent trifluoride. Polyhalide salts dissolved in non-aqueous media are also known to act as halogenating agents. For example, polybromides have been used in the bromination of jS-naphthol, styrene, aniline and phenol and in the preparation of 2,2-dibromopropanol840. !004 G. J. Perlow and M. R. Perlow, Rev. Mod. Phys. 36 (1964) 353; Chemical Effects of Nuclear Transformations, Vol. II, p. 443, Int. Atomic Energy Agency, Vienna (1965); / . Chem. Phys. 41 (1964) 1157.

1556

CHLORINE, BROMINE, IODINE AND ASTATINE: A. J. DOWNS AND C. J. ADAMS

Likewise tetrabutylammonium tetrachloroiodate(III) has been shown to be a chlorinating agent in 1,2-dichloroethane, since, upon illumination, the solution yields the (1-chlorobutyl)tributylammonium ion840. Apart from dissociation, the polyhalide anions are also susceptible to solvolysis, disproportionation, oxidation and exchange. Thus, in aqueous solution the anions all suffer some degree of hydrolysis, the tri-iodide being most resistant but such resistance diminishing as iodine is replaced by the more electronegative halogens. Further, on the evidence furnished, for example, by the voltammetric behaviour of halide-halogen or halide-interhalogen systems, unsymmetrical trihalide anions are prone to disproportionate in favour of symmetrical species (see p. 1541). The oxidation of a polyhalide anion, notably by a halogen or interhalogen molecule, leads either to a new polyhalide anion richer in the number of halogen atoms it contains or, where such a species is unstable, to the cor­ responding interhalogen compound: e.g. IC1 2 -+IC1 - > I 2 C 1 3 - 840 MCIF4+F 2 -> MF+CIF5

958

Polarographic and chronopotentiometric studies show that in solvents which stabilize the 13" ion, e.g. acetonitrile, nitromethane, acetone or 1,2-dimethoxyethane, anodic oxidation of the iodide ion occurs in two steps, viz. 31-—113-—13/212 840 However, the intermediate formation of the 13 ~ ion is not apparent in the one-step electrooxidation of iodide which is observed with solutions in water, alcohols or acetic acid. For certain redox couples implicating stable polyhalide species, e.g. I2/I3" and I3 ~/I ~, standard potentials have been determined. That exchange processes such as *I-+I3-

i a *i 3 -+i-

occur very rapidly in solution has been established by experiments using radioactive halogen isotopes1005; for the bimolecular reaction between * I _ and I3-, 127I nmr measure­ ments signify that k2 (at 27°C), ΔΗ2Χ and Δ5 2 ί are 2·2χ 109 s e c-i mol-i 1, 4-5kcal mol-i and 0·7 eu respectively (see p. 1243). Despite the fact that polyhalide anions enter into many of the reactions of halogen and interhalogen molecules and are potentially significant as intermediates, for example, in reactions involving electrophilic attack by the neutral molecules, details about specific reactions of the polyhalides, other than dissociation, are still comparatively sparse. Accordingly, a clearer definition of the chemical functions and reactivity of the anions must await the findings of more comprehensive studies. 4. BONDING I N NEUTRAL AND ANIONIC POLYHALOGEN SPECIES834,835,978,i003,i006,i007

An adequate interpretation of the electronic structure and bonding within an inter­ halogen molecule or polyhalide anion must take account of the geometry and dimensions, as well as the thermodynamic and spectroscopic properties, which have been described in 1005 M . F . A . D o v e and D . B . Sowerby, Halogen Chemistry (ed. V. G u t m a n n ) , Vol. 1, p . 4 1 , Academic Press (1967). 1006 E . H. Wiebenga and D . Kracht, Inorg. Chem. 8 (1969) 738. 1007 B . M. D e b and C. A . Coulson, / . Chem. Soc. (A) (1971) 958.

BONDING IN NEUTRAL AND ANIONIC POLYHALOGEN SPECIES

1557

the preceding sections. Several distinct approaches to the problem have been essayed. The first model to be proposed invoked essentially electrostatic interactions between ions and polarized molecules to account for the formation of polyhalide anions834»835. To describe the valency state of the central halogen atom in species such as C1F3 or IC14", a second approach assumed that electrons are promoted from the s- or /7-orbitals to the vacantrf-orbitalsof the central atom; hybridization of the valence orbitals, represented, for example, as sp*d or sp*d2, then gives rise, via the principle of maximum overlap, to localized electron-pair bonds834»835. An alternative view does not specify the involve­ ment of d-orbitals, but takes the localized electron-pairs to be a consequence of the repulsions between electrons of the same spin implicit in the Pauli exclusion principle1008. In yet another approach, delocalization of the /^-electrons of the halogen atoms via multicentre molecular orbitals is considered to be the primary stabilizing influence in the forma­ tion of a neutral or anionic polyhalogen aggregate834»835'978»1003»1006. With this premise, simple and modified Hückel procedures have been used in an attempt to explain the structures, stabilities and spectroscopic properties of polyhalogen systems1006. The most sophisticated calculations so far undertaken, those of Deb and Coulson1007, employ all the valence-shell atomic s- and p-orbitals as basis functions, in conjunction with CNDO/2 and INDO methods, to deduce ground-state properties of interhalogen molecules. The Electrostatic Model According to the electrostatic model, the formation of a trihalide anion is the outcome of ion-dipole interaction between the halide ion and a point dipole induced by it in the associated halogen molecule. Hence, the stability of the trihalide complex is evidently a function of the polarizability and radius of the constituent halogen atoms. It has been shown that the combination of ion-dipole attraction, Born repulsion and van der Waals' attraction does lead to the right order of magnitude for the energy of formation of the I 3 ~ ion835. However, the significance of the calculations is questionable in view of several empirical factors which they include and of the very crude approximation of an induced point dipole, which is certainly far from valid. Electrostatic interaction between an I - ion and an I2 molecule inevitably results in an unsymmetrical 13" ion because the two bonds are, by their nature, non-equivalent. It follows that the centrosymmetric I 3 ~ ion found, for example, in [AsPh4][l3] and [NEt4]I7 is not compatible with the electrostatic model, unless it is assumed that the interaction involves an I + ion and two I ~ ions or that the pressure of the environment induces a centrosymmetric ion. Although the model gives a satisfactory qualitative account of the structures adopted by the I 5 ~ and I 8 2 _ ions and by the anionic aggregate of [NMe^Ip, it is not directly amenable to other polyhalide anions or to neutral interhalogen molecules. It must be concluded, therefore, that a satisfactory explanation of the bonding in polyhalogen systems demands at least some degree of electrondelocalization. Localized Covalent Bonds Covalent bonding in a diatomic interhalogen molecule XY presents no problems, being analogous to that in the homonuclear molecules X2 and Y2. Similar molecularorbital schemes apply to both types of molecule, which owe their stability to the bonding capacity of the doubly occupied σ-orbital derived almost exclusively from the /?-orbitals 1008 R . j . Gillespie and R. S. Nyholm, Quart. Rev. Chem. Soc. 11 (1957) 339; R. J. Gillespie, Angew. Chem., Internat. Edn. 6 (1967) 819.

1558

CHLORINE, BROMINE, IODINE AND ASTATINE: A. J. DOWNS AND C. J. ADAMS

of the halogen atoms (see Fig. 11). In order to accommodate the additional bonds formed by one or more halogen atoms in all the other polyhalogen complexes, some form of hybridization involving */-orbitals has commonly been presumed. As the promotion of a valence electron to a d-orbital in a halogen atom requires a very large amount of energy (see Table 10), it has to be assumed that this energy can be recouped by the formation of strongly directed hybridized bonds and by the improved separation of the charge clouds of the lone-pair electrons when they go into the hybrid orbitals. Polarization of the d-orbitals of the central halogen atom by the surrounding electronegative ligands is likely to be an important factor in facilitating the participation of these orbitals in bond formation1009. The promotion of one electron in the ground-state configuration ns2np5 of the halogen atom to an nd-ovbital furnishes three unpaired electrons in the valence state of the atom, and thus provides for the formation of up to three bonds with other halogen atoms. In order to account for the observed shapes of the molecules or ions which result, the central atom is assumed to employ five sp*dhybrid orbitals; according to experi­ mental experience, the d2* orbital is thus implicated in hybridization and the orbitals are directed to the corners of a trigonal bipyramid. In a molecule such as CIF3 or BrF 3 , two of the hybrid orbitals are used to accommodate the lone-pair electrons on the central atom, while in a trihalide ion such as ICI2 ~, there are three lone-pairs to be accommodated in this way. Promotion of two electrons to «^-orbitals and the assumption of octahedrally disposed sp*d2 hybrid orbitals by the central atom are the corresponding prerequisites to the forma­ tion of interhalogen molecules of the type XY 5 or of pentahalide anions such as BrF 4 _ or ICI4 ~. ThefiveX-Y bonds of the XY 5 molecule account for five bonding-pairs of electrons, which, with a lone-pair, complete the valence shell of X, whereas in an anion of the type XY4" the six hybrid orbitals centred on X are occupied by four bonding- and two lone-pairs of electrons. Finally, the structure of IF 7 can be rationalized in terms of sp*d* hybridization, requiring the promotion of three electrons to i/-orbitals. Hence, all the valence electrons of the iodine atom are used in bond-formation. Depending on the precise proportions of s-, p- and ^/-functions which make up the hybrid orbitals, the bonds in individual polyhalogen aggregates may be equivalent or non-equivalent. Hence, it is possible qualitatively to understand the variations of interatomic distance within a particular aggregate. For instance, the disparity in bond lengths in the CIF3 or BrF3 molecule (Table 93) can be rationalized by the assumption that the heavy atom draws on some form of /?
BONDING IN NEUTRAL AND ANIONIC POLYHALOGEN SPECIES

1559

around a central nucleus, each sphere representing the orbital of a pair of electrons1010. In reality, electron clouds separated by a mean distance r do not suffer purely coulombic interactions with an associated repulsive force F = e2/r2, nor are they strictly to be treated as hard objects such that F = e2//·00; instead a force law with F = e2/rn9 where n commonly takes values between 6 and 12, is likely to give a more realistic account of the interactions. Hence, the most stable arrangement of 5 electron-pairs is at the vertices of a trigonal bipyramid for 2 < n < oo, though the square pyramid is inferior by a comparatively small margin. By the same token, 6 electron-pairs should be octahedrally disposed for 2 < n < oo. For 7 electron-pairs, however, it has been found that, as n decreases, the most stable arrange­ ment changes from 1:3:3 to 1:4:2 and finally to 1:5:1 corresponding to the pentagonal bipyramid (see Fig. 41); calculation shows that the energies of all three arrangements are

FIG. 41. Possible stereochemistries of IF7 and related species in which the valence shell of the central atom contains seven electron pairs100*.

very similar. The susceptibility of a molecule like IF 7 to facile intramolecular exchange ("pseudorotation") thus finds a ready explanation. It is also evident that the trigonalbipyramid and pentagonal-bipyramid arrangements of electron pairs present non-equivalent apical and equatorial sites. For example, each apical electron-pair in the former arrange­ ment has three nearest neighbours at 90°, while each equatorial pair has only two such nearest neighbours at 90° and two at 120°. For n ^ 4, the total force experienced by the apical pairs exceeds that experienced by the equatorial pairs, and equilibrium is reached only if the apical pairs are at a greater distance from the nucleus of the central atom than are the equatorial pairs. If it is also accepted that the charge clouds of lone-pair (LP) electrons are larger and occupy more of the "surface" of the central atom than the bonding electrons (BP), then the repulsive forces must decrease in the sequence LP-LP > LPBP > BP-BP, and it follows that the replacement of bonding-pairs by lone-pairs is attended by a decrease in the angles between the bonding-pairs. With these assumptions, it is possible to explain the geometries of C1F3 and BrF3 and of symmetrical trihalide anions. Here the lone-pairs of electrons occupy the more spacious equatorial sites in the trigonal bipyramid so as to minimize the interactions between electron-pairs mutually disposed at 90°. As with other molecules in which the central atom has 5 electron-pairs in its valenceshell, the apical bonds of CIF3 and BrF3 are longer than the equatorial bond, the presence of the lone-pairs in the equatorial belt having the effect of amplifying the disparity. Like­ wise, a simple explanation is found for the geometries of molecules of the type XY 5 and of i°i° H. A. Bent, / . Chem. Educ. 40 (1963) 446.

1560

CHLORINE, BROMINE, IODINE AND ASTATINE! A. J. DOWNS AND C. J. ADAMS

anions such as BrF4~. The enhanced repulsion between a lone-pair and adjacent electronpairs then accounts for the elongation of the basal relative to the apical bonds of XF 5 molecules and for the i/ms-arrangement of the lone-pairs in the pentahalide anions. Moreover, the different spatial demands of bonding and non-bonding electron-pairs in molecules of the types XF3 and XF 5 lead to interbond angles significantly less than the ideal values of 90° and 180°. Of the 7 electron-pairs which formally make up the valenceshell of the central atom in the anions BrF6 _ and IFa ", one is a lone-pair, and it has been argued that the most stable arrangement is the 1:3:3 structure (Fig. 41), in which the lonepair occupies the unique axial position and so has the minimum number of nearest neighbours1008. Such a structure probably represents the most stable configuration of the molecule XeF6 " 7 , but the facility with which both XeF 6 and IF 7 896 undergo intramolecular rearrangement may mean that crystal or solvation forces ultimately determine the structures adopted by BrF6 ~ and IF 6 -. Analogous to these species are the ions SeX62 " and TeX62 _ (X = Cl, Br or I), which are characterized by regular octahedral structures, although the isoelectronic complexes SbX63 - appear to vary their stereochemistry to suit the environ­ ment1011. Against the successes of those models which invoke localized covalent bonding must be set a number of weaknesses or failures. In general, the burden of assumptions denies any quantitative treatment, and the models leave unexplained some of the finer structural and energetic points. The consensus of recent opinion is that the contribution of d-orbitals is greatly exaggerated in the simple hybridization model; although the principle of electronpair repulsion does not specify whether the ^/-orbitals of the central atom are involved, its basis becomes obscure unless such participation is assumed1006. Systems poorly served by the models of localized covalent bonding include solid IC1, IBr and I2C16, the trihalide anions, and the salts [NMe4][I5], [NEt4]I7, [NMe4]I9 and Cs2l8· Thus, localized electronpairs are not easily reconciled with the shapes of the I5 _ and I 8 2 - anions or with the struc­ tural mutability of the I3- ion; more generally, there is no satisfactory basis for explaining the wide range of halogen-halogen distances within these aggregates. Delocalized Bonding To counter the shortcomings of earlier models, molecular-orbital schemes have been devised for polyhalogen systems to provide, in effect, for delocalization of the valence p-electrons of the halogen atoms engaged in bonding834'835,978,i003. The treatment typically entails, in its simplest form, the following assumptions834»835: (i) The molecular orbitals ψΗ (k = 1 . . . n) are represented by linear combinations of the valence np-orbitals φι (i = 1 . . . n) of the halogen atoms; completely neglected as bonding agents are the low-energy ns- and high-energy «^/-functions of each atom. (ii) Overlap integrals Su for neighbouring and more remote pairs of atoms are neglected. (iii) Bond or interaction integrals H{j = $φι*Ηφ}ατ (where / and j refer to atomic orbitals of different atoms) are taken to be zero, except for those orbitals of adjacent atoms which interact in the direction defined by the bond-axis, and with which a common value of H{j = β is associated. (iv) To evaluate the coulomb integrals, different approximations have been made. As a zero approximation, all the integrals a are supposed to have the same value regardless of the nature of the halogen atoms and their formal charge. In a first approximation, the nature and formal charge of the halogen atoms are acknowledged by the following rules: ion C. J. Adams and A. J. Downs, Chem. Comm. (1970) 1699.

BONDING IN NEUTRAL AND ANIONIC POLYHALOGEN SPECIES

1561

(a) For I, Br, Cl and F the values are taken to be a, a -|-0*2/2, α+0·4β and α-}-1·0β, re­ spectively, the increments being approximately proportional to the corresponding differences in the electronegativities of the atoms. (b) A formal charge g, calculated for a given atom in the zero approximation, is supposed to increase the absolute value of its coulomb integral by 0-2Qß. Because these numerical factors are relatively arbitrary, the first approximation cannot be expected to give truly quantitative results, though it should indicate the sense in which the results obtained in the zero approximation must be modified in a more realistic account. The application of this relatively naive Hiickel theory leads therefore to the molecular orbitals n

(el

where Cf is the coefficient of the ith atomic orbital in the &th molecular orbital. According to the variation principle, so that the electronic energy Ek corresponds to the eigenvalues determined by the secular equation \Hi}-EkSij\

= 0

Of the different possible configurations of a polyhalogen aggregate, that corresponding to the lowest calculated Ek should be the most stable. Thus, for a symmetrical trihalide anion, the idealized energy-level diagram takes the form illustrated in Fig. 4; if account is taken of lateral overlap of /?-orbitals to implement π-bonding, however, a more realistic diagram is that of Fig. 40. A fundamental prerequisite to such an approach is that the atoms must be sufficiently close together to allow the non-zero bond integrals Hu to assume appreciable values. For example, to form 13-, the I~ ion must first be brought close to the I 2 molecule before an appreciable delocalization energy can be realized. In this respect, there is a nice distinction between the formally similar triatomic systems H + H 2 and I - 4-I2 (see p. 1550). With H + H 2 the Born repulsion tends to keep the components well separated because the van der Waals' attraction between them is very weak; accordingly the combined action of the various forces leads to a very loosely bound complex, having a trifling delocalization energy. However, the approach of I 2 and I - is attended by relatively strong van der Waals' and electrostatic attractions, which tend to bring the species quite close together. Furthermore, the overlap of 5p-functions starts to become effective at interatomic distances relatively large compared with those appropriate to the 1 ^-functions which provide for delocalization in the system H + H2. In general, it appears that appreciable stabilization by delocalization of/7-electrons, whether in interhalogen molecules, polyhalide anions or noble-gas com­ pounds, is promoted by the longer-range forces of van der Waals' and electrostatic interactions. Because of the complexity of a system such as 13-, however, there is no pre­ dicting a priori whether the anion will be symmetrical or unsymmetrical. The simple LCAO-MO method outlined here leads to the following conclusions834»835: (a) The scheme accounts in a reasonable way for the predominance of bond angles close to 90° or 180° in polyhalogen systems. (b) Where different arrangements of the atoms are possible, the configuration with the lowest calculated energy corresponds with that observed for the polyhalogen molecule

1562

CHLORINE, BROMINE, IODINE AND ASTATINE! A. J. DOWNS AND C. J. ADAMS

or ion. For instance, in the zero approximation the delocalization energy of a trihalide anion is zero when the bond angle at the central atom is 90° but 0·83β when the angle is 180°; the reverse order of stability is found for a trihalogen cation like IC12 +. The energy of the linear IBrCl ~ ion has also been shown to be at a minimum when the iodine atom occupies the central position. Again, the V-shaped configuration emerges with the lowest energy for a homonuclear pentahalide anion like I s - , but gives place to the square-planar configuration for the ICI4" anion. However, while these calculations serve as guides to the configuration favoured by a given number of atoms, they give little hint of how the stabilities of complexes vary with the nature or number of halogen atoms: for example, the calculated delocalization energies of the molecules C1F3 and I4 make little distinction of stability. (c) For most of the known polyhalogen systems the calculations have given estimates of bond orders and of the net charges borne by the atoms. It appears that the bond orders, which fall in the range 0-6-1 -0, invariably comply with the patterns established by measured interatomic distances and stretching force constants. If the asymmetry commonly associated with trihalide anions like 13" is correctly assumed to reflect the influence of crystal forces, then adjustment of the coulomb integrals again provides qualitative agreement between the calculated bond orders and observed bond lengths. The net charges on the halogen atoms are found to vary widely, even in homonuclear aggregates. In the symmetrical I3ion, for instance, the net charge on the terminal atoms is approximately — 0-5e, whereas that on the central atom is close to zero. Analogous calculations for ICI4 ~ imply that the net charges are — 0·52β for each of the chlorine atoms and + lO7e for the central iodine atom. Although undoubtedly exaggerated by the simple model, these are in tune with the measured nqr parameters of polyhalogen compounds917»1001. (d) The calculated energies for the three possible modes of dissociation of the BrICI ~ anion show that the process BrICl~(g) -> Cl~(g) + IBr(g) offers the minimum thermodynamic barrier. This factor may well augment the decrease in lattice energy in accounting for the observation that crystalline salts of the BrICI" anion dissociate on heating to give IBr and the corresponding monochloride (see p. 1541). Qualitative agreement between theory and experiment has also been obtained for the energies of formation of diatomic interhalogen molecules in the gas phase835. Various refinements and extensions of the simple molecular-orbital approach have been essayed. Thus, a more sophisticated view of the electronic structure of polyhalogen systems must take account of (/?-/>)7r-orbitals associated with individual halogen-halogen bonds and of their capacity to mix with the (p-p)ff-orbitals when the interbond angles deviate from the limiting values of 90° or 180° " 8 . Some weight must also be given to the bonding contributions of ns~ and «^f-orbitals. Hence, the σ-bonding orbitale of the ICI4 ~ anion are undoubtedly augmented through interactions involving the 5dx*_y* orbital of the central iodine atom, while (/>->rf)7r-bonding may well provide an important mechanism for the transfer of charge between halogen atoms978. The most recent refinement of the Hückel approach1006 retains the exclusive use of valence /?-orbitals as basis functions, but takes into account (a) the electrostatic interaction of the net charges on the atoms, and (b) the change of energy of each atom, including its core electrons, as a function of its charge. The iterative procedures of this modified Hückel theory have been used to recalculate the net charges, bond orders and energy terms of polyhalogen compounds. In their general tenor, the results resemble those derived from the simple model, but the implied reduction

BONDING IN NEUTRAL AND ANIONIC POLYHALOGEN SPECIES

1563

in the net charges borne by the atoms makes for greater realism and for improved quantita­ tive agreement with structural, thermodynamic and spectroscopic findings. In a more elaborate description of the electronic properties of interhalogen molecules, Deb and Coulson1007 have considered as basis functions all the valence-shell atomic orbitals other than d. They have applied separately the approximate CNDO/2 and INDO methods in a consistent and non-empirical fashion to determine the energies of the different configurations open to individual molecules, and then used the theoretical equilibrium geometries to calculate dipole moments, orbital energies and ionization potentials, electronic transition energies, harmonic force constants and other molecular parameters. If the calculations have failed sometimes correctly to assess the relative stabilities of alternative configurations, they have enjoyed more success in explaining the equilibrium shape of a given configuration and in reproducing experimental trends in certain properties. Accordingly, the neglect ofrf-orbitalswould appear to be vindicated. Jahn-Teller and "Pseudo-Jahn-Teller" £#^1012,1013 A quite different approach, outlined in Section 1 (p. 1115), employs perturbation theory to relate the energy of a polyhalogen aggregate to the deformation of its ground-state geometry. On this basis, the structural characteristics are a function of the symmetry properties and relative energies of the ground and excited electronic states, which in turn bear upon the force field operative over the system. The possibility then exists that the structure of the aggregate may be susceptible to spontaneous distortion (a) if an orbitally degenerate ground state is implicated (the Jahn-Teller effect), or (b) if there exists an excited state ψΜ separated from the ground state φ0 by a comparatively small energy gap ΕΜ — Ε0 (the "pseudo-Jahn-Teller" effect)1012. Hence, calculations relating to the electronic and geometric structures of the molecules PF5, AsF 5 and BrF5 not only support the assumption that J-orbitals of the central atom play only a small part in the bonding, but rationalize the structural differences in terms of Jahn-Teller distortion of the trigonal-bipyramid model of BrF5 1013. Repulsions between the lone-pair and bonding-pairs of electrons in this molecule then manifest the nodal properties of the highest occupied molecular orbital. The ease with which certain polyhalogen units are deformed, whether reversibly via rapid intramolecular exchange, as with IF7, or irreversibly under the influence of their environ­ ment, as with 13 -, may be associated with the proximity of ground and excited electronic states (e.g. < 4 eV for I3-). In general, the mode of deformation is then appointed by the symmetry properties of the excited state. Thus, the excited electronic state Ση + of the centrosymmetric I3- ion dictates the off-centre displacement of the central atom, which is accommodated by the au+ vibration v3. The anomalously small force constant associated with the deformation thusfindsan explanation (see p. 1553), while the additional stabilization liable to accrue through the mixing-in of the excited state allowed by the deformation ultimately favours a permanently unsymmetrical I 3 ~ unit. Similar arguments apply to the anions BrF6 ~ and IF 6 -, which may be expected to share the stereochemicalflexibilityof the isoelectronic species XeF 6 997 , and consequently to be unusually responsive to the influence of their surroundings.

1012 L . S. Bartell, / . Chem. Educ. 45 (1968) 754; R. G. Pearson, / . Amer. Chem. Soc. 91 (1969) 4947. 1013 R . s . Berry, M. Tamres, C. J. Ballhausen and H. Johansen, Acta Chem. Scand. 22 (1968) 231.

1564

CHLORINE, BROMINE, IODINE AND ASTATINE: A. J. DOWNS AND C. J. ADAMS

D. ORGANIC POLYVALENT HALOGEN DERIVATIVES 1. INTRODUCTION

Despite its title, the present section is restricted mainly to the chemistry of organic compounds containing iodine(III) and iodine(V). The obvious limitations of space and context preclude any treatment perse of conventional univalent organo-halogen compounds, e.g. the alkyl, aryl and acyl halides. Moreover, by comparison with the iodine systems, the array of known organic polyvalent chlorine and bromine compounds is very small, consisting almost entirely of some chloronium and bromonium derivatives which are discussed below alongside their iodine(III) analogues. No iodine(VII) compounds contain­ ing carbon-iodine bonds have as yet been prepared, but some aromatic perchloryl derivatives are known. Organic polyvalent iodine compounds were first described by Willgerodt1014 and have been the subject of two detailed reviews1015»1016. They may be classified according to the oxidation state of the iodine atom and the number of carbon-iodine bonds on the same iodine atom. Iodine(III) Compounds (a) with one carbon-iodine bond, viz. iodo disalts RIX2 and iodoso compounds RIO; (b) with two carbon-iodine bonds on the same iodine atom, viz. iodonium compounds RR'IX, including iodolium compounds (1); (c) with three carbon-iodine bonds on the same iodine atom, e.g. PI13I.

(1)

Iodine(V) Compounds (a) with one carbon-iodine bond, e.g. iodoxy compounds R I 0 2 and "iodoso disalts" RIOX 2 ; (b) with two carbon-iodine bonds on the same iodine atom, e.g. iodosyl salts R 2 IOX. In each class the majority of stable compounds contain aromatic organic residues; Scheme 19 illustrates the major reactions by means of which the various aryl derivatives may be interconverted. Kinetically stable non-aromatic derivatives are known in which the organic groups are resistant to nucleophilic attack at the 1-carbon atom, and the current development of imaginative synthetic procedures should increase the number of these compounds in the near future. ion c . Willgerodt, / . prakt. Chem. 33 (1886) 154. 1015 p. M. Beringer and E. M. Gindler, Iodine Abstr. and Revs. 3 (1956). 1016 D . F. Banks, Chem. Rev. 66 (1966) 243.

1565

ORGANIC POLYVALENT HALOGEN DERIVATIVES PhlF,

Ph3P

PhIO, CHECLi Ph(CH=C)ICl IOC Ac,0.

Ph.II

Ph,IO(OH)

PhJO(OAc)

SCHEME 19. Interconversions of aromatic polyvalent iodine compounds.

Although cyclic chloronium, bromonium and iodonium ions (2) have long been postu­ lated as intermediates in organic substitution and addition reactions1017, stable organochloronium and organobromonium compounds were first isolated only in 19521018. The chemistry of these derivatives has been developed largely by the Russian school of Nesmeyanov.

(X=Cl,Br,I)

V)

The problems of structure and bonding presented by organic derivatives of polyvalent halogens are closely akin to those of the interhalogens and polyhalide ions discussed above in Section 4C. Chemically the compounds are noticeable for the ease with which one or more bonds to the positive halogen may be fractured; whether this cleavage is heterolytic or homolytic in nature depends on the system in question. There is insufficient evidence TABLE 101.

Compound

C-I bond length

(A)

/MC6H4NO QH5ICI2 (C6H5)2IC1 (C6H5)2lBF4

2-00* 200±005 2-08* 202 ±003

/>-ClC6H4I02

1-93*

a

CARBON-IODINE BOND LENGTHS

Reference M. J. Webster, / . Chem. Soc. (1956) 2841 E. M. Archer and T. G. D. van Schalkwyk, Acta Cryst. 6 (1953) 88 T. L. Khotsyanova, Doklad. Akad. Nauk, S.S.S.R., 110 (1956) 71 Yu. T. Struchkov and T. L. Khotsyanova, Bull. Acad. Sei., U.S.S.R. (1960) 771 E. M. Archer, Acta Cryst. 1 (1948) 64

No limits of error quoted.

MI? For a review of this subject see: J. G. Traynham, / . Chem. Educ. 40 (1963) 392; B. Capon, Quart. Rev. Chem. Soc. 18 (1964) 45. 1018 R . B . Sandin and A. S. Hay, / . Amer. Chem. Soc. 74 (1952) 274.

1566

CHLORINE, BROMINE, IODINE AND ASTATINE: A. J. DOWNS AND C. J. ADAMS

accurately to chart the variation in the properties of the carbon-halogen bond throughout this series of compounds; such carbon-iodine bond lengths as have been determined crystallographically are listed in Table 101.

2. IODO DISALTS ArIX2

Aromatic iodo disalts ArIX2 are derivatives of the hypothetical acids ArI(OH)2; in solution they are readily and reversibly converted into the iodoso compound. ArIX2 i = ± ArlO HX

They may be prepared by two general routes, involving either oxidation of an aryl iodide or substitution of X~ into some other organo-iodine(III) derivative. Combination of ICI3 with an organometallic compound is not a good method for the synthesis of ArICl2, since only in special cases does the reaction stop at this stage; iodonium compounds typically constitute the major product. Iodo dichlorides, probably the most important members of this class, are best obtained by passing dry chlorine into a chloroform solution of the parent iodo compound1019. ArI+Cl2^ArICl2 The reaction is reversible and the dichlorides dissociate more or less readily on heating or on dissolution in polar media; electrophilic substituents in the aromatic group facilitate this decomposition, as indexed by the variation both of the dissociation constant in solution1020 and of the thermal stability of the solids1021. On standing, the compounds evolve hydrogen chloride in a photo-catalysed reaction, e.g. hv

C 6 H 5 IC1 2 -+/>-ClC 6 H 4 I+HCl

Iodobenzene dichloride is used as a mild chlorinating agent in organic synthesis. Chlorination of olefinic functions1022 proceeds preferentially by a free radical mechanism [Scheme 20(a)] giving predominantly the ira/is-product; in the presence of water or of radical inhibitors, reaction proceeds more slowly by an ionic mechanism [20(b)]. By contrast, iodobenzene dichloride acts in aromatic chlorination reactions subject to metal halide catalysis merely as a source of molecular chlorine1023, the rate-determining step being dissociation of the dichloride. Other disalts crystallize (i) following dissolution of the iodoso compound in an acid [e.g. ArIF2, ArI(OAc)2], (ii) after addition of acid to a solution of ArlO or ArICl2 in acetic acid [e.g. ArI(ON0 2 ) 2 ] 1024 , or (iii) after treatment of ArICl2 with a silver salt in acetonitrile [e.g. ArI(OCOCF3)2]1025. About the white unstable solids ArIF2 little is known; iodo dibromides, diiodides and dipseudohalides are apparently unstable at 1019 H. J. Lucas and E. R. Kennedy, Organic Syntheses, Collective Vol. Ill (editor-in-chief E. C. Horning), p. 482, Wiley, New York (1955). 1020 R . M . Keefer and L. J. Andrews, / . Amer. Chem. Soc. 82 (1960) 4547. 1021 D . A. Bekoe and R. Hulme, Nature, 177 (1956) 1230. 1022 D . D . Tanner and G. C. Gidley, / . Org. Chem. 33 (1968) 38. 1023 R . M . Keefer and L. J. Andrews, / . Amer. Chem. Soc. 79 (1957) 4348. 1024 R . B . Sandin, Chem. Rev. 32 (1943) 249.

IODO DISALTS ArIX 2

1567

ambient temperatures1025. MeCN

PhICl 2 +2AgX

-40°c

> Phi+2AgCl+X2

(X = Br, I, CN, NCO)

Derivatives of weak oxyacids are unstable with respect to formation of the iodoso compound1025. MeCN

PhICl 2 +2AgN0 2

-40°O

> PhIO+2AgCl+N203

Iodo dicarboxylates, which may also be produced directly from Arl and percarboxylic acids, are used as oxidizing agents in organic synthesis. Like lead tetraacetate, ArI(OAc)2 acts to cleave 1,2-glycols or to substitute methyl groups into aromatic molecules. (a) Radical mechanism: Phiei 2

>=<

——

+ ci

Phici +- ci Cl

—

^

Cl

>

(x

Cl

+

PhIC1

2

S

-

)—(.

+

PhlCl

(b) Ionic mechanism:

H cl

+

_

■"-

Phi

. ^

+

CM V*""^

^~Λ

+ Cl

C1

/

SCHEME 20. Mechanisms of olefinic chlorination using iodobenzene dichloride.

There is no evidence that iodo disalts undergo ionic dissociation in solution. The PI1ICI2 molecule revealed in the crystal structure of this compound1026 is T-shaped, contain­ ing a linear symmetrical Cl-I-Cl group arranged almost perpendicular to the plane of the ring; the structure may be described as a trigonal bipyramid with the phenyl group and two lone pairs in the equatorial positions and chlorine atoms at the apices (Fig. 42). There are, however, significant intermolecular I-Cl contacts (3-40 A), close to those found in ICl3,SbCl5 (p. 1351). In /7-ClC6H4ICl2 the iodine atom enjoys a similar environment, although the molecule as a whole is planar1021. The Mössbauer1027 and 35Clnqr1028 spectra of PWCI2 are consistent with a bonding scheme involving only/j-functions of the iodine atom. Similarly the influence of substituents at the aromatic ring on the stability and spectra of iodo disalts is explicable solely in terms of inductive effects, and there is no cause to invoke resonance interactions involving the iodine 5d-orbitals and the aromatic π-cloud. 1025 N . W. Alcock and T. C. Waddington, / . Chem. Soc. (1963) 4103. 1026 E . M. Archer and T. G. D . van Schalkwyk, Ada Cryst. 6 (1953) 88. 1027 B . S. Ehrlich and M. Kaplan, / . Chem. Phys. 54 (1971) 612. 1028 j . c . Evans and G. Y.-S. Lo, / . Phys. Chem. 71 (1967) 2730.

1568

CHLORINE, BROMINE, IODINE AND ASTATINE: A. J. DOWNS AND C. J. ADAMS

Regarding non-aromatic derivatives, simple alkyl iodo dichlorides are formed by reac­ tion between iodoalkanes and chlorine, but are unstable even at low temperature1029; iodomethane dichlori^e, the most stable member of the series, decomposes at — 30°C to

^ ψ

245X

©>

© F I G . 42. Structure of iodobenzene dichloride.

methyl chloride and iodine monochloride. On the other hand, bridgehead iodides (e.g. 1-norbornyl)1030, α-iodomethylsulphones RSO2CH2I1031, and simple substituted iodoalkenes (e.g. CICH^CHI)1029 afford on chlorination reasonably stable dichlorides which resemble their aryl counterparts. Perfluoroalkyl iodo difluorides R,IF 2 have also been described1032 as low-melting solids obtained by fluorinating R/I (Rf = C2F5, n-C3F7, n-C4F9) with fluorine or halogen fluorides; the compounds are monomeric in the liquid phase. 3. IODOSO COMPOUNDS

Aryl iodoso compounds ArlO are pale yellow solids which disproportionate slowly at room temperature, rapidly at 100°C. 2ArIO-*ArI+ArI0 2

They may decompose explosively on melting. Methods for their preparation include: (i) the basic hydrolysis of iodo disalts in aqueous solution; (ii) direct oxidation of Arl, e.g. with ozone, fuming nitric acid or potassium perman­ ganate; (iii) the action of iodosyl sulphate, (IO) 2 S0 4 (p. 1352), on ArH in the absence of excess sulphuric acid. Chemically, iodoso compounds behave as anhydrides of the hypothetical dibasic acids ArI(OH)2. Exchange studies using H 2 0 1 8 disclose that iodosobenzene undergoes 30% exchange in neutral solution and 60% in basic solution1033. C6H5IO+H2Oi8 ^ C6H5I^ 1 ^C6H5IOi8+H20 L X)i8HJ The protonated species C6H5IOH + is believed to be the active intermediate in the oxidation of 1,2-glycols to 1,2-diketones, which is readily effected by iodosobenzene. 1029 j . Thiele a n d W . Peter, Annalen, 369 (1909) 119; J. Thiele a n d H . H a a k h , ibid. p . 131. 1030 j . B . Dence and J. D . Roberts, / . Org. Chem. 33 (1968) 1251. 1031 O. Exner, Coll Czech. Chem. Comm. 24 (1959) 3567. 1032 c . S. Rondestvedt, j u n . , / . Amer. Chem. Soc. 91 (1969) 3054. 10331. p . Gragerov a n d A . F . Levit, / . Gen. Chem. (U.S.S.R.) 33 (1963) 536.

DIORGANOHALONIUM COMPOUNDS

1569

The iodoso function is highly susceptible to nucleophilic attack, which may be intra­ molecular if the aromatic ring contains a nucleophilic group in the oriAo-position. Thus the oxidation of 2-iodoisophthalic acid produces the lactone (3) rather than 2-iodosoisophthalic acid1034. The reaction of equivalent amounts of ArI0 2 and ArlO in the presence of silver oxide, ArI02+ArIO+AgOH -* Ar2I+OH" +AgI03 probably proceeds via attack of ArK>3H- on the iodoso derivative. O—I—O

(3)

Structurally little is known about iodoso compounds, since no definitive studies have been made. The Mössbauer spectrum of solid iodosobenzene1027 is consistent with a bent C-I-0 unit (Z. C-I-0 ca. 90°), the iodine atom drawing almost exclusively on its /?-orbitals for the purposes of bonding. Intense absorptions at 600-700 cm - 1 in the infrared spectra of iodoso compounds are attributed to fundamentals in which I-O stretching plays the major part1035. 4. D I O R G A N O H A L O N I U M

COMPOUNDS

Diaryliodonium compounds ArAr'IX are known containing a wide variety of anionic groups (X) and both carbocyclic and heterocyclic aromatic residues. The anions may be readily interchanged by standard methods, and two general routes are available for syn­ thesis of the cations. (i) The acid-catalysed condensation of iodoso compounds with aromatic hydrocarbons gives good yields of iodonium compounds. Ar'IO+ArH

HX

> ArAr'IX+ H z O

Intramolecular coupling processes are particularly fruitful, e.g.1036

HS0 4

Oxidation of aryl iodides to iodoso compounds in the presence of hydrocarbons affords an opportunity for this reaction, while symmetric iodonium bisulphates Ar2I + HS0 4 ~ result from the treatment of ArH with iodosyl sulphate in the presence of excess sulphuric acid. 1034 w . C . Agosta, Tetrahedron Letters, (1965) 2681. 1035 c . Furlani and G. Sartori, Ann. Chim. (Rome), 47 (1957) 124; M. V. King, / . Org. Chem.26(\96\) 3323. 1036 j . Collette, D . McGreer, R . Crawford, F . C h u b b a n d R . B . Sandin, / . Amer. Chem. Soc. 78 (1956) 3819.

1570

CHLORINE, BROMINE, IODINE AND ASTATINE:

A. J. DOWNS AND C. J, ADAMS

(ii) Iodine trichloride combines directly with many organometallic reagents (e.g. Ar2Hg, ArSnCl3> to give Ar2ICl. A neat variation of this method exploits the instability of triorganoiodine(III) derivatives1037, C1CH=CHIC12

ArLl

> [ClCH=CHIAr 2 ] - * Ar 2 ICl+CH=CH

toluene

Diarylchloronium, diarylbromonium and diaryliodonium salts have been isolated following the thermolysis of diazonium compounds in the presence of the appropriate aryl halide1038. PhN 2 + BF 4 - +PhX -> Ph 2 X + BF 4 ~ + N 2 (X = Cl, Br or I)

Only exceptionally do yields exceed 10% of the diarylhalonium salt, although intramolecular coupling is more successful1018.

of^o — oOo HSO"

( X - C l , Br, I;

HSO"

Y « S , SO, SO,!

1039

The bromine(III) ylide (4) has also been reported

.

Bu l

o=

0- B r + -0" o H

Bul

Pr 1

(4)

Diarylhalonium compounds are potent arylating agents, undergoing thermal or photo­ chemical cleavage under mild conditions1038. PhOH CSHN P h 2 0 + Phi + H C 1 Ph,ICI 5 » C'H N , Ph7Se P h 3 S e + B F ^ * - r _ Ph 2 lBl· -±-U

Ph 4 Pb

-

Ph2ISbC:i4

Pb

Sb

Ph2BrBF4

Tl

^

Λ^-Ph tf^H + Phi β

{ ^ N + — P h BF 4 ~

Ph,Tl + BF 4 "

*- Ph2SbCI + Ph 2 SbCl, +Ph 3 SbCU

Reactions of diarylhalonium halides proceed largely via homolysis of the halogen-carbon bond and follow radical pathways. By contrast, the intervention of Ar + cations is impli­ cated in the arylation effected by the tetrafluoroboronates on the double salts with metal halides. Chlorine and bromine derivatives react more readily than their iodine analogues, although their relative inaccessibility precludes their extensive use in this respect. 1037 p. M. Beringer and R. A. Nathan, / . Org. Chem. 34 (1969) 685. 1038 A . N. Nesmeyanov, L. G. Makarova and T. P. Tolstaya, Tetrahedron, 1 (1957) 145. 1039 w . H. Pirkle and G. F. Koser, / . Amer. Chem. Soc. 90 (1968) 3598.

DIORGANOHALONIUM COMPOUNDS

1571

In the structural chemistry of halonium compounds Ar2XY, attention has centred on the nature of the X-Y bond. In polar solvents ionic dissociation is favoured, but the in­ fluence of other solvents has not been adequately studied. Evidence for the solid state rests on two-dimensional crystal-structure determinations of three diphenyliodonium salts. Discrete angular PI12I+ cations are recognizable in PI12IBF41040. PI12I units of very similar geometry feature in diphenyliodonium chloride and iodide1041, although in these halides dimeric molecules are found (Fig. 43), which are obviously closely related to the I2C16 dimer (Section 4C). Ph2ClX and Pr^BrX (X = Cl, I, BF4) are isomorphous with their iodonium analogues1042. In these systems spectroscopic techniques may not be a sensitive index to structure1043.

FIG. 43. Structure of diphenyliodonium chloride. Although some iodonium compounds have been prepared containing one non-aromatic group, e.g. Ph(CH=C)ICl, true alkylhalonium derivatives are not well known. On the basis of kinetic and stereochemical evidence, cyclic halonium ions (2) have long been indicted as intermediates in organic substitution and addition reactions. More recently, spectroscopic Br+

Br 3 "

(5)

studies of solutions of aliphatic halides in "super-acids", such as SbF 5 -S0 2 mixtures, have revealed the formation of cyclic ions [(CH2)nX] + (n = 2, 4; X = Cl, Br, I) by abstraction of X - from X(CH2)«X, and of open-chain dialkylhalonium ions RR'X + (R, R' = Me, Et, 1040 Yu. T. Struchkov and T. L. Khotsyanova, Bull. Acad. Sci.9 U.S.S.R. (1960) 771. 1041 T . L. Khotsyanova, Doklad. Akad. Nauk, S.S.S.R. 110 (1956) 71. 1042 T . L. Khotsyanova and Yu. T. Struchkov, Soviet Phys. Crystallog. 2 (1957) 378. 1043 A . N . Nesmeyanov, L. M. fipshtein, L. S. Isaeva, T . P. Tolstaya and L. A . Kazitsyna, Bull. Acad. Sei., U.S.S.R. (1964) 574.

1572

CHLORINE, BROMINE, IODINE AND ASTATINE: A. J. DOWNS AND C. J. ADAMS

Pr; X = Cl, Br, I) from the corresponding alkyl halide(s)1044; the open-chain ions possibly play a part in Friedel-Crafts alkylation and isomerization reactions1045. Curiously, the only stable alkyl derivative so far isolated is (5), obtained as a crystalline solid during the bromination of adamantylideneadamantane in CCI4 solution1046. Attempts to synthesize dialkyliodonium salts R2IX by the direct methods of iodoso coupling or organometallic exchange have not been successful, even for cyclopropyl, 1-norbornyl or 9-triptycyl derivatives1030. 5. TRIARYLIODINE COMPOUNDS

These compounds, which result from combination in an inert solvent of an organolithium compound with an iodonium salt, are nearly all unstable towards spontaneous decomposition into free radicals. Two relatively stable derivatives are known, namely triphenyliodine (from PhLi and Ph 2 I + I~) and (6) (from PhLi and dibenziodolium iodide)ioi6,i047.

6 (6)

6. IODINE(V) COMPOUNDS

Iodoxy compounds are formed in the slow disproportionation of iodoso derivatives: 2ArIO->ArI+ArI02 the simultaneously produced iodoarene is readily separated by steam distillation. Oxida­ tion of iodoso to iodoxy compounds is also effected by aqueous hypochlorite solution or by Caro's acid. The white solids ArI02 are subject to detonation on impact or on heating. They are slightly soluble in water, where they behave as weak monobasic acids. ArI02+2H20 ^ H 3 0 + + ArI02(OH)The pKa of iodoxybenzene is 10-41048. On the other hand, amphoteric character is indicated by protonation in acidic media. In cold alkaline solution the conjugate base PhI02(OH) - is quantitatively and irrever­ sibly converted to diphenyliodyl hydroxide, a strong oxidizing agent. 2PhI02(OH)- -* Ph2IO+OH- +I0 3 " +OH~ Addition of acetic acid or trifluoroacetic acid produces acetate or trifluoroacetate salts from which other diphenyliodyl compounds, Ph 2 IO + X~ (X = F, Cl, Br, IO3), may be 1044 G . A. Olah and J. M. Bollinger, / . Amer. Chem. Soc. 89 (1967) 4744; ibid. 90 (1968) 947; G. A. Olah, J. M. Bollinger and J. Brinich, ibid. 90 (1968) 2587, 6988. 1045 G. A. Olah and J. R. DeMember, / . Amer. Chem. Soc. 91 (1969) 2113. 1046 j . Strating, J. H. Wieringa and H. Wynberg, Chem. Comm. (1969) 907. 1047 K . Clauss, Ber. 88 (1955) 268. 10481. Masson, E. Race and F . E. Pounder, / . Chem. Soc. (1935) 1669

ASTATINE: INTRODUCTION

1573

1049

obtained metathetically . The infrared spectra of the carboxylates indicate an ionic formulation, viz. [Ph2IO]+[RC02] ~. Reaction of hot 40% hydrofluoric acid with an iodoxy compound gives an aryliodoso difluoridei050. ArI02+2HF -> ArIOF2+H20 With sulphur tetrafluoride, iodoxybenzene is converted to PhIF41()5i, while the related perfluoroalkyl derivatives R/IF4 (Rf = C2F5, C3F7 or n-C4F9) result from fluorination of RrI with excess of C1F3, BrF3 or BrF5H>32. To judge by their i^F nmr spectra, the tetrafluoroiodo- molecules have square-pyramidal structures, related to that of IF5, with apical organic groups. In keeping with the behaviour of inorganic iodates and iodine pentoxide, the infrared spectra of iodoxy, iodoso and iodyl derivatives are marked by strong absorptions in the region 650-800cm -1 , presumably arising from I-O bond-stretching vibrations. The two I-O linkages in/7-ClC6H4I02 are unequal in length (1-60 and 1-65 A; Z. O-I-O = 103°; Z. O-I-C = 94-5°), and probably entail some sort of multiple bonding; short intermolecular I · · · O interactions (2-87 A) must also represent a major binding force within the crystal1052. The Mössbauer spectrum of iodoxybenzene itself indicates some involvement of the iodine 5^-orbital in the bonding102?. 7. PERCHLORYL COMPOUNDS

Perchloryl aromatic compounds1053 are prepared from perchlorylfluorideby two routes: QL. CIO3F Q a Q j + UF

They are stable compounds which can be distilled in steam; the — CIO3 group resembles —N0 2 in the way that it deactivates the aromatic nucleus. Hydrolysis in concentrated alkali replaces — CIO3 by —OH. Hydrogenation on a platinum catalyst proceeds with rupture of the C-Cl bond, but the CIO3 group is unaffected by reductants such as LiAlH4, SnCl2 or zinc and hydrochloric acid. 5. ASTATINE 105 «- 10 « 5.1. INTRODUCTION

The element of atomic number 85, the ultimate member of the Halogen Group, was 1049 F . M. Beringer and P. Bodlaender, / . Org. Chem. 33 (1968) 2981. 1050 R . p . Weinland a n d W . Stille, Ber. 3 4 (1901) 2631. 1051 L . M . Yagupol'skii, V. V. Lyalin, V. V. O r d a a n d L . A . Alekseeva, / . Gen. Chem. (U.S.S.R.) 3 8 (1968) 2714. 1052 E . M . Archer, Acta Cryst. 1 (1948) 64 1053 v . M . Khutoretskii, L . V. Okhlobystina a n d A . A . Fainzil'berg, Russ. Chem. Rev. 3 6 (1967) 1 4 5 ; J. F . Gall, Kirk-Othmer Encyclopedia of Chemical Technology, 2 n d edn., Vol. 9, p . 598. Interscience, N e w Y o r k (1966). 10 54 A . G, Maddock, Mellows Comprehensive Treatise on Inorganic and Theoretical Chemistry, Supple­ ment I I , Part I, p . 1064. L o n g m a n s , G r e e n a n d C o . , L o n d o n (1956). 1055 K . W . Bagnall, Chemistry of the Rare Radioelements, p . 97. Butterworths, L o n d o n (1957). 1056 E . K. Hyde, / . Chem. Educ. 36 (1959) 15. 1057 E . A n d e r s , Ann. Rev. Nucl. Sei. 9 (1959) 2 0 3 . 1058 A . H . W . Aten, j u n . , Adv. Inorg. Chem. Radiochem. 6 (1964) 207. 1059 v . D . Nefedov, Y u . V. Norseev, M . A . T o r o p o v a a n d V. A . K h a l k i n , Russ. Chem. Rev. 37 (1968) 87. 1060 D . R . Corson, K . R . MacKenzie a n d E . Segre, Phys. Rev. 57 (1940) 4 5 9 , 1 0 8 7 ; Nature, 159 (1947) 24. 1061 (a) E . H . A p p e l m a n , U . S . A t o m i c Energy Commission Reports U C R L - 9 0 2 5 a n d N A S - N S 3012 (1960); (b) E . H . Appelman, MTP International Review of Science: Inorganic Chemistry Series One, Vol. 3 (ed. V. G u t m a n n ) , p . 181, Butterworths a n d University P a r k Press (1972).

1574

CHLORINE, BROMINE, IODINE AND ASTATINE: A. J. DOWNS AND C. J. ADAMS

named astatine by Corson, MacKenzie and Segre, who in some experiments reported in 1940 were the first to prepare and identify unambiguously an isotope of this element and to explore some of its chemical properties1060. The name, which was derived from the Greek word αστατοσ meaning "unstable", emphasizes the instability of all the known isotopes of the element with respect to radioactive decay. For a number of reasons astatine is one of the most difficult elements to investigate from a chemical point of view. (i) Thefirstisotopes to be discovered, 211At and 210At, with half-lives of 7-2 and 8-3 hr, respectively, are still, and will probably remain, the longest-lived isotopes. The short halflife of even the most suitable isotopes makes it virtually impossible to work with the element in a laboratory which does not itself have the facilities to produce the isotopes. Accordingly the larger part of our present knowledge concerning astatine depends on the activities of only a small number of investigators; following the pioneering work of Corson, MacKenzie and Segr£, the elegant studies of Appelman1061, also at the Berkeley Radiation Laboratory, have been pre-eminent in bringing some qualitative and semi-quantitative sense to the seemingly erratic behaviour of the element. (ii) Astatine being chemically available only as a result of nuclear synthesis, its prepara­ tion is made more difficult because none of its isotopes is accessible by neutron irradiation. Instead of the more conventional and convenient nuclear reactor, an accelerator is required to produce the high-energy a, 12C or 14 N particles that are the necessary reagents for the synthesis of the longest-lived astatine isotopes. (iii) Chemical operations with milligram amounts of these isotopes would be quite impossible, even if such quantities were available, the specific activity of 210At being, for example, about 2000 C/mg, a level which would elicit enormous heating and radiation effects. For this reason, all information on the chemistry of astatine has necessarily been derived from radiochemical studies of trace concentrations, typically 1 0 _ n to 1 0 ~ 1 5 M ; the most concentrated aqueous solution of astatine recorded by Appelman1061 was no more than 10 ~8 M. The specific activity of more concentrated solutions is such as to induce radiolysis of the water, leading to the formation of hydrogen peroxide and oxygen and either masking or influencing the behaviour of the astatine. (iv) At the dilutions which must be employed to explore the chemistry of astatine, con­ centration makes a major contribution to the thermodynamic behaviour, and the extra­ polation of results to deduce macroscopic behaviour is correspondingly uncertain. These uncertainties will be most prominent in the event of dissociation or ionization or if the reaction involves a change in the number of astatine-containing species. The need for caution is underlined by observations implying that the chemistry of iodine on a tracer scale occasionally differs notably from its macrochemistry1062: not only do new ionic species appear at low concentrations, but the situation is further complicated by reactions with ever-present impurities, by adsorption effects, or by radiocolloid formation. (v) As with other tracers, the chemistry of astatine has been studied mainly by extraction and by coprecipitation, neither of which provides results easy to interpret. Although astatine should follow iodine in many of its chemical reactions, so that iodine is the obvious carrier element (cf. the use of caesium as a carrier for francium), the analogy between the two elements often leaves much to be desired, and in many situations the astatine tracer

1062 M. Kahn and A. C. Wahl, / . Chem. Phys. 21 (1953) 1185; M. L. Good and R. R. Edwards, / . Inorg. Nuclear Chem. 2 (1956) 196.

HISTORY, DISCOVERY AND NATURAL OCCURRENCE OF ASTATINE

1575

does not follow the iodine carrier in the way one would expect. For example, ions or mole­ cules of astatine compounds often do not fit well into the host lattice of the corresponding iodine compound; astatine compounds are often more strongly adsorbed than the corre­ sponding iodine compounds; again, following redox reactions with certain reagents, it is possible that iodine and astatine appear in different oxidation states. Any one of these com­ plications may lead to apparent inconsistencies in the behaviours of the astatine tracer and the iodine carrier. Although investigations of the chemistry of astatine were initiated more than 30 years ago, much of the behaviour of the element is still in doubt, and even some of the conclusions drawn from the known experimental results may well turn out to be erroneous on closer acquaintance. In particular, very little is yet known about the physical chemistry of astatine, though interesting estimates have been made of the lattice energies of alkali-metal astatides, the dissociation energy of the At 2 molecule and the electron affinity of the At atom106*, and, on the basis of chemical studies, an approximate potential diagram has been deduced for the chemistry of the element in aqueous solution1064. From its position in the Periodic Table, one may expect the element to resemble iodine fairly closely, though some signifi­ cant differences should also exist. For example, the + 7 and — 1 oxidation states of astatine should be somewhat less stable, and the electropositive character of the lower oxidation states more pronounced, than in the case of iodine. In evidence of these differences, perastatates emerge, reputedly, only under the most energetic conditions of oxidation, the astatide ion appears to be a stronger aqueous reducing agent than the iodide ion, and polyhalide complexes such as AtCl2~ have significantly greater formation constants than their iodine counterparts. The incipient metallic character of solid iodine should be more highly developed in solid astatine, in keeping with the transition from semi­ conductor to metal in the neighbouring elements tellurium and polonium. As yet the At2 molecule has not been characterized, and it is perhaps suggestive that, whereas the molecules HAt, MeAt, Atl, AtBr and AtCl have been identified by mass spectroscopy, the At2 molecule has eluded detection1065. Inevitably more polarizable and probably having lower ionization potentials, astatine atoms and molecules are likely to experience relatively stronger intermolecular or charge-transfer interactions than do the atoms and molecules of the lighter halogens. To this extent the At2 molecule is likely to be at once a better donor and a better acceptor than the other halogen molecules; it would therefore be interesting to know whether the molecule retains its identity in the solid element or whether, in the interests of long-range stability, astatine, like polonium, adopts a threedimensional lattice, wherein atoms rather than molecules are the only recognizable subunits. It has even been suggested, albeit on the basis of very limited evidence, that in its properties astatine shows a closer resemblance to polonium and bismuth, its horizontal neighbours in the Periodic Table, than to iodine. 5.2.

HISTORY, DISCOVERY AND NATURAL

OCCURRENCE"^-1059

The discovery and interpretation of atomic numbers by Moseley and the establishment of the physical basis of the Periodic Classification of the elements by Bohr in the early 1920's indicated that a fifth halogen might exist1066. Consideration of the properties of the neighbouring elements suggested that it would probably prove to be radioactive, though no 1063 M . F . C . Ladd and W . H . Lee, / . Inorg. Nuclear Chem. 2 0 (1961) 163. 1064 E . H . Appelman, / . Amer. Chem. Soc. 8 3 (1961) 8 0 5 . 1065 E . H . Appelman, E . N . Sloth and M . H . Studier, Inorg. Chem. 5 (1966) 766. 1066 E . Wagner, Z. Elektrochem. 16 (1920) 2 6 0 .

1576

CHLORINE, BROMINE, IODINE AND ASTATINE: A. J. DOWNS AND C. J. ADAMS

basis for the estimation of its radioactive properties was known at that time. The subse­ quent search for the element followed two principal lines: one involved chemical studies, usually combined with a sensitive physicochemical method of detection, especially X-ray emission spectroscopy; the other depended on radioactive studies. It was realized that, unless the element possessed a stable or very long-lived isotope, it must occur in association with one of the three natural decay series, so that attention was focused on uranium and thorium minerals, though some investigators, heedless of the likely radioactive character of the element, explored the possibility of its natural occurrence in macroscopic quantities. Several reports of the discovery of element 85 appeared as a result of these investigations. Names such as Alabamine1067, Helvetium1068, Anglo-Helvetium1069, Leptine1070 and Daccine1071 were variously assigned to the element, as were definite chemical properties. However, most of these early claims were rejected by a critical reinvestigation of the methods of separation and identification1072. Hevesy1073 and Andersen1074 were unable to detect astatine in nature, though the methods used by the latter have been shown to be inapplicable to the separation of astatine, and the results are therefore inconclusive1060. On investigating the possible radiochemical derivation of astatine, Hahn1075 observed that the currently recognized radioactive decay processes could produce isotopes of element 85 only by α-decay of an isotope of the unknown element 87 or by ß-decay of an isotope of polonium. The possible genesis of element 85 by α-decay of element 87, francium, linked the searches for these two elements, whose histories have several features in common. From the vantage point of our present, more sophisticated knowledge of nuclear stabili­ ties, we know that the search for stable or moderately long-lived isotopes of astatine was foredoomed to failure. However, with the development of a sufficiently large cyclotron, Corson, MacKenzie and Segr£ had access to beams of α-particles of sufficiently high energy to surmount the coulomb barrier of the heavy elements and so reach an energy range where the cross-section for α-induced nuclear reaction becomes appreciable. By bombardment of a bismuth target with high-energy α-particles they succeeded in producing astatine according to the nuclear scheme given below1060: 209B 83 B

1067 p . Allison, E . R . Bishop a n d A . L . Sommer, / . Amer. Chem. Soc. 54 (1932) 616. 1068 w . Minder, Helv. Phys. Acta, 13 (1940) 144. 1069 A . Leigh-Smith a n d W . Minder, Nature, 150 (1942) 767. 1070 c . W . Martin, Nature, 151 (1943) 309. 1071 R . D e , Separate, Bani Press, Dacca (1937). 1072 B . Karlik and T. Bernert, Sitzungsber. Akad. Wiss. Wien, Abt. Πα, 151 (1942) 255; Naturwiss. 30 (1942) 685; B . Karlik, Monatsh. 77 (1947) 348. 1073 G . v. Hevesy a n d R . Hobbie, Z. anorg. Chem. 208 (1932) 107. 1074 E . B . Andersen, K. Danske Vidensk. Selsk. 16 (1938) N o . 5 , 1 . 1075 o . H a h n , Naturwiss. 14 (1926) 158.

HISTORY, DISCOVERY AND NATURAL OCCURRENCE OF ASTATINE

1577

The 211 At isotope thereby produced exhibited α-activity of two kinds decaying with a halflife of about 7 hr; the α-particles of lower energy {ca. 40%) correspond to direct a-decay of the isotope to 207Bi, while those of higher energy (ca. 60%) correspond to orbital electron capture in the first place to produce a polonium isotope 211Po, whose α-decay to 2o?Pb has a very short half-life. The α-activity characteristic of 211At could be separated from the target by distillation and was also shown to be separable from mercury, lead, thallium, bismuth and polonium. The possibility of fission was excluded by both chemical and phys­ ical considerations. Although the chemical properties described in the first publication were subsequently proved to be rather inaccurate, that the α-activity was correctly attrib­ uted to the isotope 211At is now beyond dispute. Astatine can be said to occur in nature only in the widest sense of the term, being a short-lived, rare branch product in the decay of the uranium and thorium radioactive series1056»1076. In accordance with the genetic relationship between polonium and astatine, Karlik and Bernert1077 and later Walen1078 observed, in studying the radioactivity of the daughter products of radon, that about 0Ό2% of Ra-A (210Po) disintegrates by ß-decay to 218 At with a half-life of about 2 sec. Following this discovery of j8-branching in decay of Ra-A, which served as an impetus to the study of other natural radioactive series, Karlik and Bernert1079 found that Ac-A (215Po) also undergoes )3-decay to the extent of 5 x 10~4% to form 215At, which undergoes α-decay with a half-life of about 10 ~4 sec, and which is Atomic number

Actinium (4n+3) series

92

7-13 X 108y

Pa231(Pa) 4

3-48 xl0 y.

91

*Λ

90

Th227(RdAc) y

%/r

w

Ra 2 2 3 (AcX) / t - 2 % ll-68d /

88

J &»>

87

W 21m / -3 :tinon)/4xl0 '

86

At 219

85

0-9m

84

Ρ θ
83

2· 16m

o?

8Z

*™

T h 2 3 1( U Y )

AC227(AC)

y

89

y ^

Bi2l,(AcC) fc

215

P0 (AcA) / 9X Pp'"(AcA) 7% 1-83 x 10-^s / ^

M

Bi' 8m

Pb 207 (AcD) / „ 0 / P b 2 1 , ( A c B ) STABLE / ° - 3 2 / 3 6 . l m

TI20W")

81

4-79m

124

128

132

136

140

144

A-Z

FIG. 44.. Actinium (4n+3) series showing the relationships of the rare branch products 223ρΓ} 2i^At and 2i5ßi. [Reproduced with permission from E. K. Hyde, / . Chem. Educ. 36 (1959) 16.] 1076 E . K. Hyde and A. Ghiorso, Phys. Rev. 90 (1953) 267. 1077 B . Karlik and T. Bernert, Naturwiss. 32 (1944) 44. 1078 R . j . Walen, / . Phys. Radium, 10 (1949) 95. 1079 B . Karlik and T. Bernert, Z. Physik, 123 (1944) 51.

1578

CHLORINE, BROMINE, IODINE AND ASTATINE: A. J. DOWNS AND C. J. ADAMS

also produced in the successive α-disintegration of 227Pa1080. On the basis of the regularities in the a- and jS-decays in the natural radioactive series, it was suggested that, as a result of rare instances of α-decay, 227Ac forms a minor decay chain parallel to the main series, which avoids element 85 (see Fig. 44)107°. In verification of this suggestion, a careful scrutiny of the actinium (4n + 3) series led to the isolation from the 227Ac source of 223Fr (half-life 21 min) and hence of an α-emitting 219At isotope (half-life 0-9 min) formed by a-branching to the extent of 4 x 10 _3 %. 219At was identified by chemical means, this being the only case of chemical isolation of astatine from a natural source1076. The situation with respect to the natural occurrence of astatine may be summarized as follows: (i) minute amounts of 218At (ti = 2 sec) and possibly of 215At (t± = 10 ~4 sec) formed in uranium ores as a result of rare ß-branching of the A-products; (ii) minute amounts of 217At (ti = 0-03 sec) from the steady production of very small amounts of 237 Np and 2 3 3 U in uranium ores through natural neutron processes; (iii) minute amounts of 219 At (/$ = 0-9 min) present in 235 U ore samples because of a very minor α-branching of 223 Fr, which itself is a rare product. 219At is the longest-lived isotope of astatine likely to be discovered in natural sources. Estimates suggest that the outermost mile of the earth's crust contains no more than about 70 mg of astatine, as compared with 24-5 g of francium or the relatively abundant polonium (4000 tons) and actinium (11,300 tons)1081. Astatine thus appears to be the rarest element in nature. 5 . 3 . I S O T O P E S OF ASTATINE1056-1059

Some 21 isotopes of astatine have now been more or less well defined. The half-lives, modes of decay and decay energies are listed in Table 102, together with details of the nuclear reactions whereby the isotopes have been obtained and identified. Isotopes possessing 126 neutrons or less (e.g. 210At and 211At) are produced by bombarding bismuth with aparticles having energies in excess of 21 MeV. More recently some of the lighter isotopes of astatine have been identified in nuclear reactions involving the bombardment of gold targets with accelerated 12C, 14 N or 1 6 0 nuclei. Most of the isotopes having more than 126 neutrons are descendants of the neutron-deficient protactinium isotopes prepared in the high-energy bombardment of thorium1082. The heaviest isotope for which data are available is 219At, the daughter of 223 Fr; heavier isotopes are certainly ß ~ unstable with halflives of less than a minute. Almost all the known isotopes of astatine are also formed by fission of thorium and uranium by high-energy protons1083»1084. The half-lives and decay energies of the isotopes both for α-emission and for electron capture are of great interest in relation to the closed configurations of 82 protons and 126 neutrons, which exercise a major influence on the stability of the heaviest nuclei1056. Figure 45 indicates the general trends in nuclear stability in the region of the elements following lead in the form of an idealized sketch of the mass-energy surface. In general, the surface takes the form of a trough sloping upwards in the direction of higher masses, and the bottom of the trough defines the region of 0-stability. The steepness of the slope determines the 1080 w . W. Meinke, A. Ghiorso and G. T. Seaborg, Phys. Rev. 75 (1949) 314. 10811. Asimov, / . Chem. Educ. 30 (1953) 616. 1082 w . W. Meinke, A . Ghiorso and G. T. Seaborg, Phys. Rev. 81 (1951) 782. 1083 F . F . Momyer, jun. and E . K . Hyde, / . Inorg. Nuclear Chem. 1 (1955) 274. 1084 M . Lefort, G. Simonoff and X . Tarrago, Bull. Soc. chim. France, (1960) 1726; B . N . Belyaev, Y u n g - Y u Wang, E . N . Sinotova, L . N e m e t and V. A . Khalkin, Radiokhimiya, 2 (1960) 6 0 3 .

210At

209At

208At

207At

206At

205At

204At

203At

202At

1

204po

205Po

206Po

207Po 203ßi 212Fr

209Rn 213Fr 210Rn

8-9-9-3 min 26-2 min

29-5-32-8 min

1-8 hr 1-6 hr 5-5 hr 8-3 hr

5-95 5-90

5-70

5-76 5-65 5-65 5-519(32%) 5-437(31%) 5-355(37%)

a(4-52%),EC(95-5%) Δ -11

a(18%),EC(82%) Δ -13

a(~88%),EC(~12%) Δ -12

a(~10%),EC(~90%) Δ -13-41

a(0-5%),EC(>99%) Δ -12

a(~5%),EC(~95%) Δ -12-89

a(0-17%),EC(>99%) Δ-1212

207R n 211Fr

210p O 206ßi

209Po 205ßi

208Po

203Po

7-4 min

609

a(14%),EC(86%) Δ -11

209Fr

202Po

3 0 min

612(64%) 6-23 (36%)

a(12%),EC(88%) Δ -10

205Fr

20ipo

201At

1-5 min

a

6-35

Parent

a, EC

Daughter

Genetic relationships

200p o

Half-life 0-9 min

Energy of a-radiations (MeV) 6-42 (-60%) 647 (-40%)

Observed mode of decay Mass excess (Δ MeV)

200At

Isotope

TABLE 102. ISOTOPES OF ASTATINE*

209Bi-a-3/i

209Bi-a-4 w

209ßi-a-5/i

209Bi-a-6n 197AU-14N-4//

197AU-12C-3/I 197AU-14N-5/Z 209Bi-a-7/i

197AU-12C-4/I 209Bi-a-8« 197AU-14N-6/I

197AU-12C-5/? 209Bi-a-9/i

197AU-12C-6/I

197AU-12C-7/I

197Au-12C-8/i

197AU-12C-9/I

Principal means of production

9-2

213At

a

a(>99%),j3-(0-l%) Δ 811

α(~97%),0-(~3%) Δ10-5

218At

219At

219Rn 215ßi

214ßi

218Po

223Fr

213Bi 22iFr

212ßi

2HBi

207ßi 21ip0

Parent

Unknown β~ emitters with high decay energies and very short half-lives

0-9 min

1-5-2-0 sec

6-70 (94%) 6-65(6%)

6-28

00323 sec

220p r

2i9Fr 215p 0

2i8Fr

2"Rn

Daughter

Genetic relationships

Descendant 227Ac natural source

Daughter 21 spo

Descendant 225Ac

Descendant 228pa

Descendant 227pa

Descendant 226pa

Descendant 225pa

209Bi-a-«

209Bi-a-/i

209Bi-a-2«

Principal means of production

a = ^He; n = neutron; EC = orbital electron capture. The means of production by nuclear reactions are represented as: target element-projectile-outgoing particle. a C . M. Lederer, J. M. Hollander and I. Perlman, Table of Isotopes, 6th edn. (1968); E. K. Hyde, / . Chem. Educ. 36 (1959) 15; V. D. Nefedov, Yu. V. Norseev, M. A. Toropova and V. A. Khalkin, Russ. Chem. Rev. 37 (1968) 87. b G. H. Fuller and V. W. Cohen, Nuclear Data Tables, 5A (1969) 459.

>219At

a Δ4-38

217At

Δ2-25 7-07

~3xl0~4sec

7-80

216At

~ 10-4 sec

801

a

215At

Δ -1-25

~ 2 x 10"6 sec

8-78

a Δ -3-42

< 1 sec

2HAt

Δ -6-5

a

212m At

0 1 2 sec

0-30 sec

7-60(20%) 7-66(80%) 7-82(80%) 7-88(20%)

a

Δ -8-64 a Δ -8-42

2l2At

2"At

7-21 hr

Half-life

5-868

Energy of a-radiations (MeV)

a(40-9%),EC(59-l%) Δ -11-64

Observed mode of decay Mass excess (Δ MeV)

/ = 9/2 b

Isotope

TABLE 102 (cont.)

ISOTOPES OF ASTATINE

1581

amount of energy available for α-decay. In the region of lead and bismuth the surface is relatively flat so that the decay energy for α-emission is very low. However, immediately after lead, in the region appropriate to isotopes of francium and astatine, the surface rises steeply, with the result that all the isotopes have very short lives. With the move to still

FIG. 45. Mass-energy surface for the heaviest elements. Atomic masses displayed on an idealized mass-energy surface to show general trends in nuclear instability. Vertical height is proportional to mass. Coordinates in the base plane are the atomic number Z and the mass number A. [Reproduced with permission from E. K. Hyde, J. Chem. Educ. 36 (1959) 15.]

heavier elements, the surface levels off somewhat, and it is in this region that the long-lived isotopes 232Th, 235U a n d 238U are found. The general features of nuclear binding which underlie the gross changes on this energy surface are rather well correlated by the indepen­ dent particle or shell model of the nucleus: as nuclei are built up by the addition of neutrons and protons, there occur favoured closed shells of neutrons and protons in a manner reminiscent of the occurrence of especially stable electron configurations in the building up of the chemical elements. 2°spb is particularly tightly bound because it contains such closed configurations or "magic numbers" of 82 protons and 126 neutrons. The additional nucleons which must be added to make the nuclides immediately succeeding lead are not as firmly bound, a feature reflected in the steepness of the mass-energy surface in the region of element 85. Though all of the isotopes of astatine with less than 126 neutrons (mass number < 211) undergo α-deeay, the half-life of this process is sometimes quite long: e.g. 5-5 years for 210At.

1582

CHLORINE, BROMINE, IODINE AND ASTATINE: A. J. DOWNS AND C. J. ADAMS

However, all such isotopes lie well to the neutron-deficient side of ß-stability and the decay energies for electron capture are several MeV. By contrast, isotopes of astatine having more than 126 neutrons are characterized by marked instability with respect to α-decay; as with francium, the discontinuity in atomic masses which occurs at 126 neutrons causes a related discontinuity in the α-decay energies. Instead of decreasing as the mass number increases in accordance with the usual pattern, these energies first decrease and then rise abruptly to a maximum once the 126-neutron shell has been completed, corresponding to the sharp decrease in binding energy of the nucleons. The heaviest astatine isotopes show a modest increase in stability with respect to α-decay, but susceptibility to ß-decay finally denies the existence of long-lived species with mass numbers in excess of 211. The ^-stability of astatine has been explored by the method of closed decay-energy cycles, which suggest that all isotopes of astatine are unstable with respect to ß-decay processes1085. The 210At isotope (t$ = 8-3 hr) is the longest-lived form of astatine at present known, 209 At and 211At having comparable lifetimes (t± = 5-5 and 7-2 hr, respectively). There remains the possibility that an isomer of one of the known isotopes will ultimately be dis­ covered with an appreciably longer half-life. One isomer 2 i 2m At has been identified with some assurance; others, less convincing, have been reported. The consensus of opinion is that 210At is the longest-lived form of astatine likely ever to be encountered and that chemical studies of astatine are likely always to be circumscribed by the constraints and ambiguities of tracer methods. 5.4.

P R E P A R A T I O N , S E P A R A T I O N A N D E S T I M A T I O N OF ASTATINE"^-io59,i06i

The chemical properties of astatine are most conveniently explored using the relatively long-lived isotopes with mass numbers 209, 210 and 211; although not quite the longestlived, 211At is, from the point of view of the chemist, the most widely favoured. All three isotopes are prepared by the α-bombardment of bismuth. The threshold α-energy for astatine production is about 21 MeV, and the cross-section for the formation of 211At rises to about 0-9 barn at 26 MeV; if the α-particles are accelerated to more than 29 MeV, the threshold of the (a, 3«) reaction is passed and 210At is produced, while 209At is formed with a fairly large cross-section by α-particles of 60 MeV. To produce 211At of the highest radiochemical purity, therefore, the particle energy should not exceed 29 MeV. Astatine is the main product of the α-bombardment of bismuth; the primary reaction does not give significant yields of other elements. The bismuth is irradiated either as the metal or as the oxide. The metal is normally applied to a support of gold, silver or aluminium by fusion or by sputtering; the oxide is most readily introduced into holes drilled in a metal slab. To avoid loss of astatine by volatilization during the irradiation, the target must be cooled efficiently. Separation of the astatine from the target is normally effected either (i) by distillation or (ii) by "wet" extraction methods. (i) To separate astatine from metallic bismuth, advantage is taken of the relatively high volatility of the astatine. Distillation is typically performed in a current of nitrogen (at a pressure of a few mm Hg) or in vacuo, the target being heated to 300-600°C. The astatine is condensed on a suitable cold surface, offered either by a glass cold finger or by a cooled platinum disc. Further purification is achieved by redistillation. The astatine is loss p . Everling, L. A. König, J. H. E. Mattauch and A. H. Wapstra, Nucl. Phys. 18 (1960) 529; A. H. Wapstra, ibid. 28 (1961) 29.

PREPARATION, SEPARATION AND ESTIMATION OF ASTATINE

1583

finally washed from the glass or metal surface on which it has been collected with a suitable aqueous solution, e.g. dilute nitric acid. If the distillation is carried out at too high a tem­ perature, the astatine is contaminated with polonium and possibly even with weighable quantities of bismuth, and tends to form radiocolloids which may seriously interfere with many experiments. The proportion of astatine effectively transferred in this way is generally low (5-30%) and irreproducible, probably because much of it is lost by adsorption or by chemical reactions on the walls of the vessels used or on other materials inevitably present. (ii) Wet methods for the isolation of astatine can be employed either for the preparation of carrier-free solutions or of solutions in an excess of iodine. In a typical procedure the irradiated bismuth or bismuth oxide is dissolved in perchloric acid containing a little iodine as a carrier for the astatine. The bismuth and some of the polonium are then precipitated as phosphates. For many purposes this liquid can be used as an aqueous solution of Atl, though it still contains some polonium, which may give rise to difficulties in certain situa­ tions. Purification can be achieved by extracting the I 2 containing Atl into carbon tetrachloride or chloroform, whence it may be re-extracted into a reducing aqueous phase. Alternatively, the astatine, initially in a hydrochloric acid solution in the presence of iron(II) salts, is extracted into di-isopropyl ether. This method has the advantage that, after the addition of a suitable quantity of tributyl phosphate, polonium and bismuth can be quan­ titatively returned to an aqueous phase containing nitric and hydrochloric acids. By the use of such wet methods losses of astatine have been restricted so as not to exceed 5%. Methods have also been devised for separating astatine from lead, bismuth, thorium and polonium following the fission of thorium under the action of high-energy protons1059. These depend upon the dissolution of the target in acid, followed by selective redox and precipitation reactions, together with solvent-extraction or Chromatographie techniques. The preparation of astatine for estimation is largely dependent on the composition of the solution and the method of analysis. If only volatile materials are present in appreciable quantities, and if one is dealing with aqueous solutions, the liquid may be evaporated to dryness on silver or platinum. However, quantitative deposition of astatine on evaporation is difficult to achieve from organic systems. Alternatively astatine may be precipitated from aqueous solutions on a silver foil after the fashion of the technique used to study polonium. Another approach is to coprecipitate the astatine with other materials: astatine(0) is coprecipitated quantitatively with the sulphides of bismuth, mercury, silver and antimony, rather less completely with HgO, Fe(OH)3, Al(OH)3 and La(OH)3. Tellurium-precipita­ tion carries astatine satisfactorily when tellurous acid is reduced with sulphur dioxide in a hydrochloric acid medium. In the presence of iodine, astatine is best isolated by the preci­ pitation of an iodide from a reducing solution. For this purpose the precipitation of palladium iodide is reported to be more reliable than that of silver iodide. The characteristic α-spectrum of a 211At sample serves as a distinctive autograph of the isotope; the 40 : 60 ratio of the 5-87 MeV α-particles of 211At and the 7-43 MeV a-particles of 211Po, and the controlling half-life of 7-2 hr for both, furnish conclusive proof of the identity and purity of a stock solution. This spectrum can be obtained with a gridded ion chamber whose output is subjected to linear amplification and pulse height analysis. If the radiochemical purity of the sample is assured, it is possible to follow the fate of astatine through tracer experiments simply by employing, for example, a liquid scintillation counter. If the purity is in question, there is the danger that 2 iop 0 may interfere with such measure­ ments. The α-activity may be measured either on a very thin or on a very thick sample; the former arrangement is more accurate for counting but the preparation is more adventitious,

C.I.C. VOL II— CCC

1584

CHLORINE, BROMINE, IODINE AND ASTATINE: A. J. DOWNS AND C. J. ADAMS

whereas the latter arrangement is generally quite easy to prepare but the assay may produce considerable errors. However, numerous investigators prefer to determine 211At by counting the 80 keV K-X-rays originating in the electron-capture branch of the astatine decay. For example, Appelman has analysed astatine solutions by dilution to a standard volume and by counting the 80 keV X-rays with a counter containing a well-type crystal of sodium iodide. The α-emission of 210At being insignificant, this nuclide is estimated by the X-rays or y-rays produced by decay; likewise 209At decays to the extent of only ca. 5% by α-emission and is most conveniently determined by X-ray or y-ray counting. For chemical purposes it does not, as a rule, matter very much if the astatine activity observed is due partly, say, to 210At and partly to 211At, though the X- and y-radiations emitted by 209At, 210 At and 211At differ widely in energy, and discrimination can be effected by a crystal with good resolution. Various methods have been described whereby absolute determinations of the activity of astatine can be carried out, though α-standardization of 211At by liquid scintillation counting appears to be most reliable. Autoradiography of astatine has also been used successfully in biological work, the activity measured being the α-emission of 211At. 5.5. PHYSICAL PROPERTIES OF ASTATINE AND ITS COMPOUNDS

Since the half-lives of known astatine isotopes are so short as to preclude the isolation of the element or its compounds in macroscopic quantities, few extranuclear physical prop­ erties of astatine-containing systems can be measured. Most of the properties described have been interpolated or extrapolated by various theoretical or empirical treatments from the data available for neighbouring or homologous elements; some of these properties are listed in Table 103. For example, a first ionization potential of ca. 9-5 eV has been esti­ mated1086, while extrapolation of the almost linear relationship between I\ for O, S, Se and Te and I2 for F, Cl, Br and I implies that I2 = 17-0 eV (rather than 18-2 eV as suggested elsewhere1086). Extrapolation of the vibrational frequencies ωβ of the familiar diatomic halogen molecules points to a value of ωβ of ca. 160 cm - 1 for the At 2 molecule. Use of the empirical relationships log we = a—Z>logn2/ logD = c+i/log/

(n = principal quantum number of the valence shell; a, b9 c and d = constants or near-constants) between ωβ and / and D, the ionization potential and dissociation energy, respectively, of the At 2 molecule then leads to the estimates of /and D given in the table1087. These results, combined with a probable value of ca. 20 kcal mol _1 for AZ/>[At2(g)], imply that astatine should be better disposed than iodine to form cations of the type At + or At2 + as chemically recognizable entities. As evaluated by several investigators1088, the conventional ionic radius of the At" ion is close to 2-3 Ä. Radius-ratio considerations suggest that CsAt is therefore likely to have a caesium chloride-type structure, Li At a zinc blende lattice and the remaining alkali-metal astatides to be isostructural with sodium chloride; together with compressibilities and other factors estimated by reference to the known alkali halides, these assumptions have been 1086 w . Finkelnburg and F . Stern, Phys. Rev. 77 (1950) 303. 1087 R . w . Kiser, / . Chem. Phys. 33 (1960) 1265. loss w . H. Zachariasen, The Actiniae Elements (ed. G. T. Seaborg and J. J. Katz), p. 775, McGraw-Hill, New York (1954); L. Genov, / . Gen. Chem. (U.S.S.R.) 29 (1959) 683; G. A. Krestov, Radiokhimiya, 4 (1962) 690.

PHYSICAL PROPERTIES OF ASTATINE AND ITS COMPOUNDS

1585

TABLE 103. ESTIMATED PHYSICAL PROPERTIES OF ASTATINE AND ITS DERIVATIVES

Value

Property Atomic properties Atomic number Electronic configuration of neutral At atom in its ground state Ionization potentials, I\

h

Electron affinity at 298°K Electronegativity Absorption spectrum of atomic astatine Properties of elementary astatine Dissociation energy of At2 molecule,

Reference

85

[Xe]4/i45i/i06i26/75 ( 2 ^3/2)

J

eV 9-5 170 3-1

kcal 220 392 71

2-4

2 lines observed

^(At 2 )298°K Δ#?.298'κίΑ«8)]

eV 1-2 10

Ionization potential, /(At2> Vibrational frequency, ωβ Melting point Boiling point

eV kcal 8-3 191 160 cm"! ~300°C ~335°C

Internuclear distance, r e (At-At)

Properties of the astatide ion Ionic radius of At" ion

2-9 Ä

kcal 27-7 24

-AH?.29e-K[At-(g>]0 -AG2 y d [At-]at298 K

2-3 Ä eV kcal 20 47 64 kcal

-Atf h 0 y d [At-]at298°K

66 kcal

Reduction potential, £: 0 (iAt 2 /At")

+0-3V

a Estimated from /i(Po) b* Based on Mulliken scale of electro­ negativity c

d d,e Estimated by ex­ trapolation d d f f g b e, estimated from ionic radius e, estimated from ionic radius h

a W. Finkelnburg and F . Stern, Phys. Rev. 11 (1950) 303. * M. F . C. Ladd and W. H. Lee, / . Inorg. Nuclear Chem. 20 (1961) 163; * also involves estimate of AHf[At2(g)]. c R. McLaughlin, / . Opt. Soc. Amer. 54 (1964) 965. d R. W. Kiser, / . Chem. Phys. 33 (1960) 1265. e See for example A. G. Sharpe, Halogen Chemistry (ed. V. Gutmann), Vol. 1, pp. 10-12, Academic Press (1967); D. A. Johnson, Some Thermodynamic Aspects of Inorganic Chemistry, pp. 70-71,101-105, Cambridge (1968). f R. E. Honig and D . A. Kramer, Vapour Pressure Curves of the Elements, Sheet C, RCA Laboratories (1969). * W. H. Zachariasen, The Actiniae Elements (ed. G. T. Seaborg and J. J. Katz), p. 775, McGraw-Hill, New York (1954); L. Genov, / . Gen. Chem. (U.S.S.R.) 29 (1959) 683; G. A. Krestov, Radiokhimiya, 4 (1962) 690. h E. H. Appelman, / . Amer. Chem. Soc. 83 (1961) 805.

shown1063 to furnish the following lattice energies (£/0>kcal mol - 1 ): LiAt, 172; NaAt, 157; KAt, 147; RbAt, 142; CsAt, 140. Heats of formation of these astatides have also been computed by an extrapolation technique, whence AH/[At ~(g)] has been evaluated. On the

1586

CHLORINE, BROMINE, IODINE AND ASTATINE: A. J. DOWNS AND C. J. ADAMS

basis of this and the estimated values of D(At2) and A///[At2(g)], the electron affinity of astatine emerges as ca. 71 kcalmol - 1 . Other derivations of thermodynamic parameters for astatine and some of its compounds have been given1059, while the ionic radius of A t _ has been applied here to the estimation of absolute values of AG° and ΔΗ° for the hydration of the ion (see Table 103). The bulk of the At" ion clearly implies that the lattice and hydration energies of ionic astatides are likely to be significantly less than those of corresponding chlorides, bromides or iodides. Accordingly the At ~ ion, even more than the I ~ ion, is liable to be unstable (with respect to redox processes) in solids containing metal ions in relatively high charge states, e.g. F e 3 + or Cu 2+ , but stable in the presence of metal ions in relatively low charge states, e.g. Cu + or Eu 2 + (seep. 1119). Likewise the solubilities of ionic astatides should follow a pattern similar to that of the corresponding iodides, though for cations of a given charge the free energy of solution should reach a maximum (corresponding to low solubility) at a larger cation radius. With B-metal ions like Ag + or Hg 2 + , however, charge-transfer and electron delocalization are likely to exercise an even greater influence in solid astatide deriva­ tives than in the corresponding iodides. In all probability the energies of covalent bonds engaging a halogen to a more electropositive atom follow the sequence F > Cl > Br > I > At, though with respect to a more electronegative partner astatine may well form the strongest bonds. 5.6. C H E M I C A L P R O P E R T I E S OF ASTATINEi056-i059,i06i

As already indicated, information about the chemical properties of astatine is available almost exclusively from tracer experiments, which are notoriously liable to misinterpreta­ tion. The following factors are typical of those which must be weighed in any attempt to interpret the behaviour of astatine under these conditions: (i) the very low concentration of the species under investigation, (ii) adsorption phenomena, (iii) interference from im­ purities, (iv) formation of radiocolloids, (v) photosensitivity of reactions, and (vi) differences in behaviour between astatine and the carrier element iodine. Notwithstanding these difficulties, at least four oxidation states of astatine are known at present: a - 1 state that coprecipitates with Agl or Pdl 2 ; a zero state which can be extracted into organic solvents; a + 5 state that is carried by insoluble iodates; and an intermediate state ( + 1 or +3?) that is neither extracted into organic solvents nor carried from solution by insoluble iodates. The redox equilibria among these oxidation states have been studied qualitatively in acidic solution using reversible redox couples and approaching the equilibria from both sides1064. On the basis of these chemical tests, the reduction potential scheme shown in Fig. 46 has been tentatively proposed; the behaviour of iodine and astatine is compared in the accompanying oxidation state diagram. Figure 47 summarizes the principal features of the aqueous chemistry of astatine that appear to have been established beyond the realm of dispute. Oxidation State 0 When astatine is isolated from a bismuth target by volatilization, it shows many prop­ erties expected of the zero oxidation state. It is volatile from glass surfaces at room tem­ perature, but is adsorbed strongly by metals such as silver, gold and platinum and weakly by others, e.g. nickel and copper. Adsorption by silver remains high at 325°C indicating the formation of a stable compound. Zerovalent astatine is soluble in nitric acid and is

CHEMICAL PROPERTIES OF ASTATINE

1587

commonly used in this form for tracer studies. Evaporation of the solution causes some volatilization of the astatine. The retention of the halogen upon various surfaces during this (a) Standard potentials E° in volts H 5 AtO ö ( ?)

>+1·6

> At03

+1-5

> HOAt( ?)

+1-0

► At(0)

+0·3

► At"

(b) Oxidation state diagram

11 10

9 8

nE° volts

5

4

3

2

1

0

Fro. 46. Suggested oxidation states and approximate reduction potentials of astatine in aqueous solution, a H + = 1*0.

evaporation depends upon the nature of the surface, the composition of the aqueous solu­ tion and the conditions of evaporation, but the volatility losses are not as high as those for similar solutions containing iodine. There exists a considerable fund of coprecipitation data relevant to the aqueous chemistry of the lowest oxidation states of astatine. Thus, astatine is carried quantitatively by insoluble sulphides precipitated from strong hydro­ chloric acid and by elemental tellurium deposited by the action of sulphur dioxide, zinc or

1588

CHLORINE, BROMINE, IODINE AND ASTATINE: A. J. DOWNS AND C. J. ADAMS Not extracted into CCl 4 but quantitatively carried on AglOj, B a f l O ^ , Pb(I0 3 ) 2 and La(OH) 3

S 2 0 8 ,Ce(IV), I0 4 " or NaBi0 3

VO, 2 + Fe 3 + (in light)

Not extracted into CCL

At(0) or Atl

VOTFe (in the dark)

Extracted into CC14

4

nor coprecipitated with insoluble iodates

Z n , S 0 2 + H + , SO* •f O H " Fe(CN)*~* or As(III)** At" I 2 ,Fe(CN) 6 3 ",t As(V)tfor dilute HNO,

Coprecipitates with Agl

AtX"

AtX

Extracted intoCCl •4

Not extracted into CCL but extracted into di-isopropyl ether

Conditions [ F e ( C N ) 6 4 " ] » [ F e ( C N ) 3 ] (E < + 0 · 4 V) * Low acidity and ionic strength, [As(III)] » [As(V)] (E < + 0 · 3 V) ** Low acidity and ionic strength, t 1:1 Fe(CN) 3~/ Fe(CN) 6 4 ~ mixture at p H < 3 ( Ε > + 0 · 4 V) f t p H < 4 (E>+0-3V)

FIG. 47. Aqueous reactions of astatine.

other strong reducing agents, but only partially by insoluble hydroxides. It is also deposited from dilute nitric acid solutions on such surfaces as copper, bismuth and silver. In this respect it closely resembles the neighbouring element polonium. The zero oxidation state of astatine resembles that of iodine in the freedom with which it is extracted from aqueous media into organic solvents, which include carbon tetrachloride, benzene and other hydrocarbons, as well as donor solvents such as ethers and tributyl phosphate. With the exception of astatine(O) and interhalogen compounds such as Atl, none of the other oxidation states is extracted into a non-polar organic solvent like carbon tetrachloride. It follows that, in principle, such solvent-extraction procedures afford an excellent means of deriving quantitative information about astatine chemistry. By observing changes in extraction coefficients with astatine concentration, with acidity, with the concentration of potential complexing agents or with oxidation-reduction conditions, it should be possible to specify the formula of the extracted species, the extent of complexing

CHEMICAL PROPERTIES OF ASTATINE

1589

and the oxidation potentials. In practice, only crude estimates of these quantities have generally been forthcoming because of the difficulty of realizing reproducible, non-drifting results. The general characteristic of solvent-extraction studies is that the nominal distribu­ tion coefficients achieved when a dilute acid solution is repeatedly put in contact with an organic solvent decrease quickly to a very low value, while the distribution coefficient for the repeated "back-washing" of the first solvent fraction rises steadily towards some limiting value of perhaps 50-100, which is not well reproduced with a different stock solution of astatine. Presumably the solvent-extraction behaviour of an element as reactive as iodine or astatine at extremely low concentrations is conditioned by adsorption on, or reaction with, impurities, colloidal particles or the surfaces in contact with the solutions. Under these handicaps only large changes in distribution coefficients are meaningful, and it has not been possible to decide on the basis of present evidence whether the species extracted by the organic phase is At 2 or At. Nevertheless, our knowledge of astatine chemistry has been substantially enlarged by qualitative studies of this sort. Thus, a decrease of the partition coefficient of astatine between the organic and aqueous layers has been observed with the following reagents™** -1059,1064: Alkali. This suggests that in aqueous alkaline solution astatine, like iodine, disproportionates to the — 1 and a higher oxidation state. Increasing the acidity of the solution restores the astatine to its initial state and it is again extracted by the organic phase. Reducing Agents such as Sulphur Dioxide, Arsenic(JII) or Metallic Zinc. Evidently astatine(O) is reduced to the astatide ion. Halide Ions, e.g. Cl~, Br~ or / - . Presumably polyhalide ions incorporating astatine are formed. The distribution coefficient diminishes markedly in the presence of a non-polar solvent like carbon tetrachloride but remains high in the presence of a donor solvent like di-isopropyl ether. Oxidizing Agents ofIntermediate Power, e.g. Cl2, Br2, Vanadate, Concentrated Nitric Acid and Fe* +. These oxidize the astatine to some oxidation state X intermediate between 0 and + 5. Very little is known about this state. Several of the reactions are further complicated, moreover, by their photosensitivity. The reaction with chlorine, bromine or iodine gives the interhalogen compounds AtX (X = Cl, Br or I) 1065 . Powerful Oxidizing Agents, e.g. S20$2~, Periodate, HOCl, Ce(JV) or Sodium Bismuthate. These convert astatine to a form which coprecipitates with insoluble iodates, presumably At03It is also significant that Fe2 + reduces higher oxidation states of astatine to astatine(0) but not to the At ~ ion1064. Distillation experiments from a variety of media (e.g. nitric, hydrochloric, perchloric and sulphuric acids) indicate that the involatile states of astatine present in most of these solutions can be reduced to volatile astatine(0) by Fe2 +. Solventextraction and coprecipitation studies show that astatine in the +X oxidation state is reduced to At(0) in the dark by the V0 2 + /V0 2 + couple (E° = +0-95 V) and by the Fe3 +/Fe2 + couple (E° = + 0-76 V) even when the ratio [Fe3 +]/[Fe2 +] is large (E = 0-89 V). Under the influence of light, however, the vanadium couple oxidizes At(0) quantitatively to At(+X) provided some iodine is present. The reaction of astatine with the ferric-ferrous couple is not sensitive to light unless [Fe3 +]/[Fe2 +] > 100 (E > +0-87 V). When the

1590

CHLORINE, BROMINE, IODINE AND ASTATINE: A. J. DOWNS AND C. J. ADAMS

ratio is greater than this, the astatine is nearly completely photo-oxidized to At( + X) whether or not iodine is present. The photochemical reactions thus appear to involve a gross shift of a chemical equilibrium under the influence of light. Such an effect is likely only if the reaction under investigation involves a species present at trace concentrations, but the primary photochemical process involves only species present at macro concentra­ tions. The information about these reactions is insufficient to justify speculation as to their mechanisms, but the results make clear the hazards of interpretation that beset tracer studies of astatine. Astatide is oxidized to At(0) by the £Ι 2 /Ι _ couple (E° = +0-54 V), by a 1:1 ferrocyanide-ferricyanide mixture at pH < 3 (E > +0-4 V) and by the As(V)/As(III) couple at pH < 4 (E > +0-3 V). On the other hand, the ferrocyanide-ferricyanide couple reduces At(0) to A t - at low acidity and ionic strength if the concentration of ferrocyanide ions is several times that of ferricyanide (E < +0-4 V), and the As(V)/As(III) couple behaves similarly at pH > 4 (E < +0-3 V). These results serve to define a value of ca. +0-3 V for £°[At(0)/At-]. In summary, astatine(O) has been characterized as volatile, soluble in organic solvents and susceptible to adsorption on various surfaces; unfortunately it is also noted for its severe irreproducibility of behaviour. Oxidation State — 1 The formation of the astatide ion by reduction of astatine(O) with a sufficiently powerful reducing agent is indicated (i) by electromigration experiments and (ii) by coprecipitation with insoluble iodides such as Agl, Til or Pdl 2 , irrespective of pH and of other anions which may be present. Unlike astatine(0), which may also be carried from solution, the astatide ion is not removed from the insoluble iodides by washing with acetone, nor is it extracted from aqueous solutions by carbon tetrachloride. Although the astatide ion probably represents the best defined chemical state of astatine, having a striking resemblance to the iodide ion, it is not retained very well by some iodide precipitates like Agl, though initially it coprecipitates quite efficiently. The ratio of the diffusion coefficients A~/^At" = 1*41 in 1% sodium chloride solution containing 1·2χ 1 0 - 3 Μ Κ Ι and 4x 10~ 3 MNa 2 SO3 has been determined1058, a value implying that the diffusion coefficients of the halide ions pass through a maximum and then decrease as a function of atomic number. The result can also be used to evaluate the mobility of the A t _ ion in aqueous solution, which likewise conforms to the pattern F~ < Cl~ <, Br _ ^ I~ > A t - . The earliest tracer studies showed signs of the formation of a volatile hydride in acid solutions containing the astatide ion. More recently the hydride has been identified by its mass spectrum1065, though whether it is produced by the hydrolysis of astatine by traces of water or by substitution reactions with organic impurities is not clear. Interhalogens and Polyhalide Ions Whereas the behaviour of astatine(0) is irreproducible, by mixing astatine(0) with other halogens or halide ions it is often possible to obtain reproducible behaviour characteristic of interhalogen species. In the presence of a large excess of elemental iodine, for example, the so-called "zero" oxidation state is represented exclusively by Atl molecules. Under these circumstances the astatine is protected against adsorption and a fairly accurate and reliable value of ca. 5-5 for the distribution ratio between carbon tetrachloride and water has been

CHEMICAL PROPERTIES OF ASTATINE

1591

obtained; for the more polar AtBr the corresponding value is 0-04. The coprecipitation of Atl with I 2 from chloroform solutions has also been studied. Mass spectrometric studies confirm that the direct addition of excess iodine, bromine or chlorine to a freshly sublimed sample of astatine converts a large part of the astatine into the corresponding monohalide106^. Although higher astatine chlorides and bromides might be expected to exist, they would probably not be formed under the rather mild halogenating conditions employed in this experiment. Curiously the action of C1F3 on astatine fails to give a volatile product. The formation of polyhalide species has been studied by means of solvent-extraction experiments. One such study showed that astatine is extracted by di-isopropyl ether from chlorine-treated hydrochloric acid solutions, affording strong evidence of the presence of a polyhalide ion, though whether this should be formulated as AtCl 2 ~ or AtCl 4 ~ could not be determined. Appelman1089 has exploited the distribution of astatine between water and carbon tetrachloride to investigate the reactions of the element with I 2 ,1 ~, IBr, Br ~ and Cl ~. Hence the species Atl, AtBr, Atl 2 ~, AtlBr ~, AtlCl ~, AtBr2 - and AtCl 2 " have been charac­ terized and the equilibrium constants interrelating them have been evaluated. These constants correlate well with analogous data relating to the lighter halogens, and it is interesting to note that astatine is always complexed a little more strongly than iodine in the corresponding circumstances, the ratios of the formation constants amounting to a factor of 2-5-8. The formation of AtCl and AtCl 2 ~ has likewise been confirmed more recently by measurements of the distribution of astatine between an aqueous oxidizing solution and metallic platinum coated with an oxide film, and stability constants for these species have thus been derived1090. The AtCl2 _ ion has also been characterized by paper electromigration experiments and is adsorbed on sulphonic-cation-exchange resins, e.g. Dowex 50. An alternative approach to the investigation of astatine-containing polyhalide species is to produce solid derivatives. For example, Csl 3 containing CsAtI2 is easily prepared by adding an excess of solid iodine to an aqueous solution of Csl containing some Atl. The thermal decomposition of CsAtI2 obeys the general rule for the decomposition of polyhalides, namely that the halogen of lowest atomic number is retained to form the involatile monohalide; hence, during the decomposition of CsAtI2 in the Csl 3 host lattice, astatine in the form of Atl is volatilized and involatile Csl remains1091. Oxidation State +X In aqueous solution there appears to exist an oxidation state intermediate between 0 and + 5. This may represent either AtO ~ or A t 0 2 -, but no convincing evidence of its chemical nature is yet available, and it must be borne in mind that so-called "At( + X)" need not represent a single species or even a single oxidation state. Indeed the nature of the state may vary with the mode of preparation. Formed by the oxidation of aqueous astatine(O) with oxidants such as bromine or dichromate or, in the presence of light, vanadate or iron(III), astatine(X) is characterized by not being extracted into carbon tetrachloride and by its failure to coprecipitate with either insoluble iodates or iodides. The oxidation state also results when astatate is reduced with chloride ions [£°(£C12/C1 -) = +1-36 V]; the periodate-iodate couple [E° — ca. + l -6 V] effects the reverse change, although the reaction is sometimes slow and incomplete at room temperature. A potential of +1-5 V is thus indicated1064 for the couple A t 0 3 ~/At(X). On the evidence of such redox behaviour, 1089 E . H . Appelman, / . Phys. Chem. 65 (1961) 325. 1090 Y U . v . Norseev and V. A . Khalkin, / . Inorg. Nuclear Chem. 30 (1968) 3239. 1091 G . A . Brinkman, J. T h . Veenboer and A . H . W. Aten, jun., Radiochim. Acta, 2 (1963) 4 8 .

1592

CHLORINE, BROMINE, IODINE AND ASTATINE: A. J. DOWNS AND C. J. ADAMS

the + X state appears to be significantly more stable than either 10 ~, which is known, or I02~~, which is not. Recent experimental results1059·10611* on the electrodeposition of astatine from oxidizing solutions imply the presence of an astatine cation having a charge of + 1 . This species— possibly At + or AtO + —is adsorbed on monofunctional sulphonic-cation-exchang eresins, on certain precipitates containing large univalent cations, e.g. caesium tungstophosphate and silver iodate, and on metallic platinum coated with an oxide film (see above). The asta­ tine deposited on platinum is held strongly on the metal surface; adsorption is increased by raising the temperature and decreased by reducing the pH; desorption is easily accomplished by anodic polarization of the platinum. Evidence of a cationic species analogous to [(C 5 H 5 N)2l] + is also forthcoming from tracer studies of the formation of [(C5H5N)2l] + X~ (X = NO3 or CIO4) in the presence of astatine1092. The ratio of astatine to iodine is higher in these derivatives than in the iodine used for the synthesis, a finding concordant with the enhanced ability of astatine to form cationic species, e.g. according to the reaction Donor species coord T Λ- coord

In all probability the free energy of this process in strongly coordinating solvents like water is most favourably disposed towards ionization in the case where X = At. Oxidation States + 5 and + 7 Treatment of astatine(O) with strong oxidizing agents converts it to a form which is not extracted into carbon tetrachloride but is quantitatively coprecipitated with insoluble iodates, e.g. AgI0 3 , Ba(I0 3 ) 2 and Pb(I0 3 ) 2 , as well as with lanthanum hydroxide. On this basis, the existence of an astatate ion, presumably formulated as A t 0 3 -, has been inferred, the negative charge on the ion being demonstrated by means of electromigration experi­ ments. Curiously it has been reported that although, in the presence of strong oxidizing agents, astatine coprecipitates well with Pb(I0 3 ) 2 , a large part of the activity due to the astatine may be eluted from the precipitate by acetone. After the most vigorous oxidation with agents such as periodate or cerium(IV) in 6 M perchloric acid, the astatine appears to be entirely in the form of A t 0 3 ~ ; precipitation of potassium periodate fails to carry any of the activity with it. Lately, however, there have been reports1093 testifying that aqueous perastatate is generated either through the agency of such potent oxidants as XeF2, K 2 S 2 O s or KOC1 or by anodic oxidation; the identity of the At04~ ion is sustained both by electrophoretic analysis and by coprecipitation with MIO4 salts (M = K or Cs) from solutions at pH 6. Organic Derivatives of Astatine1058ao59,i06ib The method of producing and studying organic derivatives of astatine which has so far proved most practicable involves the use of iodine as a carrier and investigation of the subsequent fate of the activity due to the traces of astatine. Identification of the astatine compounds is achieved most conveniently with the aid of Chromatographie techniques. However, the difficulties encountered in working with the inorganic forms of astatine are exacerbated in the study of organic derivatives of the element. Irregular extraction and 1092 j . j . c . Schats a n d A . H . W . A t e n , j u n . , / . Inorg. Nuclear Chem. 15 (1960) 197. 1093 v . A . K h a l k i n , Y u . V . Norseev, V . D . Nefedov, M . A . T o r o p o v a a n d V . I . K u z i n , Doklady Chem. 195 (1970) 8 5 5 ; G . A . N a g y , P . G r o z , V . A . H a k i n , D o K i m T u o n g a n d Y u . V . Norseev, Kozp. Fiz. Kut. Intez. Kozlem. 18 (1970) 173.

CHEMICAL PROPERTIES OF ASTATINE

1593

coprecipitation, combined with the relative instability of many astatine compounds, often make the interpretation of observations quite uncertain. If, for instance, iodoform is synthesized from iodine containing Atl, thefirstprecipitate of CHI3 shows a very appreciable activity, which disappears rapidly in the course of a series of recrystallizations. It is not clear whether this signifies that the lattice did not contain CHAtI2 at all or whether this compound, once formed, is very easily decomposed, though the first explanation seems more plausible. This situation probably explains the early rather conflicting reports made about the synthesis of organic compounds of astatine1058. A general consideration of the chemical properties of astatine suggests that the best chance for incorporation of astatine into organic molecules may exist in those cases where the astatine atom carries a partial positive charge, both in the final product and in any intermediate stage through which it passes during the reaction. In keeping with this general observation, methods have been described1059»1094 for the synthesis of astatine-labelled compounds, not only of the type RAt, but also of the types RAtCl2, R2AtCl and RAt0 2 (where R = phenyl or/7-tolyl); the details are outlined in the following scheme:

R2IC1 + KI(At)

R 2 Hg r-^-—R 2 I(At)Cl 170-190°C Cl 2 ^ R 2 M ( A t ) — ^ R l ( A t ) — - RI(At)CM lNaOC»RI(At)02 70-100°C

A variety of methods has been exploited lately to synthesize astatobenzene1095, e.g. At" Ph 2 H

Phl PhN2Cl

175°C y Ph 2 IAt

Atl

► P h i + PhAt

>PhAt

<

^PhAt

<

Atl 2 ~, I3-

130-200°C

PhNHNH 2 Phi

Reference has also been made to hot-atom reactions which apparently afford the oppor­ tunity of generating carrier-free astatobenzene: in one such procedure triphenylbismuth is irradiated with α-particles, and in another the electron-capture decay of 211Rn to 211At is engineered in the presence of benzene. Successful syntheses have been claimed, moreover, for /?-astatobenzoic and />-astatobenzenesulphonic acid1058; the latter is of interest as providing a means of incorporating astatine into different proteins, a process which can also be accomplished by means of diazo-coupling of benzidine to protein. Elsewhere there are good grounds for believing that the interaction of the At" ion, under appropriate conditions, with an alkyl iodide1096 or with ICH2COOH1097 furnishes the corresponding organo-astatine compound. Separation and characterization of organic derivatives have commonly entailed the use of thin-layer, paper or gas chromatography, the last of these techniques inviting 1094 v . D. Nefedov, Yu. V. Norseev, Kh. Savlevich, E. N. Sinotova, M. A. Toropova and V. A. Khalkin, Doklady Chem. 144 (1962) 507. J 095 G. Samson, Chem. Weekbl. 65 (1969) 27; G. Samson and A. H. W. Aten, jun., Radiochim. Acta, 13 (1970) 220; V. D. Nefedov, M. A. Toropova, V. A. Khalkin, Yu. V. Norseev and V. I. Kuzin, Radiokhimiya, 12 (1970) 194; V. I. Kuzin, V. D. Nefedov, Yu. V. Norseev, M. A. Toropova and V. A. Khalkin, Radio­ khimiya, 12 (1970) 414. 1096 G . Samson and A . H. W. Aten, jun., Radiochim. Acta, 12 (1969) 5 5 ; M. Gesheva, A . Kolachkovsky and Y u . V. Norseev, Preprint, Joint Institute for Nuclear Research, D u b n a , P6-5683 (1971). 1097 G . Samson and A . H . W. Aten, jun., Radiochim. Ada, 9 (1968) 53.

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CHLORINE, BROMINE, IODINE AND ASTATINE: A. J. DOWNS AND C. J. ADAMS

estimates of the boiling points of several volatile astatine compounds1095»1096. In aqueous solution the acid AtCH2COOH appears, on the evidence of solvent-extraction experiments, to be marginally weaker than ICH2COOH, the pKa values at 22°C being 3-70 and 3-12, respectively1097. 5.7. B I O L O G I C A L B E H A V I O U R OF ASTATINEi058,i059

Preparations of astatine for biological purposes typically involve distillation in vacuo from a bismuth target and condensation in a trap on a thin ice surface; after melting the ice, an aqueous solution of astatine in minimal volume is obtained. Solutions used for physiological studies are prepared by the addition of the appropriate amount of sodium chloride to the aqueous solution of astatine. For such investigations very small quantities of astatine (< 10 ^C) are usually employed, and to obtain precise information about the distribution of astatine, a precise analytical method of separating it from biological speci­ mens has been developed by Hamilton and his coworkers1098. The organic substance is destroyed by wet combustion with a mixture of perchloric and nitric acids. The solution is subsequently evaporated or diluted with distilled water to bring the concentration of per­ chloric acid to 3 M; the astatine is then separated from this solution for counting by coprecipitation with metallic thallium or by adsorption on a silver disc, the latter being more suitable for the serial analysis of biological materials. The distribution of astatine can also be observed in autoradiograms, in which the 211At is recognized by its a-tracks. It has been observed that, if astatine is administered as a radiocolloid, it tends to accumu­ late in the liver, a tendency common to most radiocolloids. However, if administered as a true solution, astatine, like iodine, accumulates in the thyroid gland, as was first demon­ strated in 1940 by Hamilton and Soley1098 when comparing the metabolism of radioiodine and astatine. It is not clear whether the astatine is carried by tW blood as At ~ or as At(0). For some time its chemical condition in the thyroid gland was also in doubt, but it has now been shown1098 that a large fraction of the astatine in the thyroid follows the protein fraction as it is precipitated by trichloroacetic acid or by concentrated ammonium sulphate. This suggests that astatine can take up positions in the thyroid protein similar to those normally occupied by iodine. The uptake of astatine by the thyroid gland is reduced by the administra­ tion of thiocyanate, a behaviour quite analogous to that of iodine. On the other hand, the influence of propylthiouracil enhances the uptake of astatine but reduces that of iodine. Relatively small amounts of astatine are capable of provoking profound changes in thyroid tissue without inducing any noticeable alterations of structure in the parathyroid gland or other peritrachial tissues. Astatine is superior to radioiodine for destruction of abnormal thyroid tissue because of the localized action of the α-particles, which dissipate 5-9 MeV within a range of 70microns of tissue, whereas the much weaker j8-rays of radioiodine have a maximum range of approxi­ mately 2000 microns in tissue. Detailed investigations have been made in guinea pigs, rats and monkeys1098. On this basis, the use of astatine to treat hyperthyroidism in humans has been proposed, though the hazards from deleterious side-effects have not been fully explored; reports that mammary and pituitary tumours can be induced in rats with a single injection of astatine do not breed confidence in the future of such treatment. loos J. G. Hamilton and M. H. Soley, Proc. Nat. Acad. Sei. U.S.A. 26 (1940) 483; J. G. Hamilton, Radiology, 39 (1942) 541; J. G. Hamilton, P. W. Durbin, C. W. Asling and M. E. Johnston, Proc. Intern. Conference on Peaceful Uses of Atomic Energy, Geneva, 10 (1956) 175.