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The ionization and dissociation energies of molecules and radicals
1.~ucmr.Spccmm. Radlor.Trcmricr. Vol.2,PP. 327-333. Pwmtmn
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THE IONIZATION AND DISSOCIATION ENERGIES OF MOLECULES AND RADICALS W. C. PRICE, P. V. HARRIS and T. R. PASSMORE King’s College, London. W.C.2, England Abstract-It is shown that in molecular ions two factors, (i) polarization and (ii) radical deiormation energies, assume as large a role as covalent and ionic binding do in neutral mok.culea. This is illustrated particularly with reference to species XI& and XFn, where X refers to atoms in the Grst row of the periodic table. Values are given for the ionization and dissociation energies of all these molecules or radicals and their ions. The unusual space dependence of polarization energy is reconciled with the
appearance of Rydberg spectra, the vibration frequencies, and interatomic distances in the ionized state. IN THE ground states of neutral molecules the main factors responsible for molecular formation are the covalent and ionic binding between the atoms. When a molecule is dissociated work has to be done against these forces, the nature of which and their variation with internuclear separation is fairly well understood. The covalent forces which arise from the overlap of bonding orbitals fall off very rapidly, whereas the ionic forces are coulombic and decrease more slowly with distance, a factor which considerably affects the shapes of the corresponding potential energy curves. In the case of the ions of molecules or radicals two additional major binding factors are involved which become of the same order as those already mentioned as of major importance in neutral molecules. These are (a) polarization energy and (b) bond deformation energy. The latter is also of importance for neutral molecules. A molecular ion is formed from a charged and a neutral radical and as these approach, the polarization of the neutral molecule gives rise to a polarization energy which is equal to &~e2/r”’at fairly large distances (a being the polarizability of the neutral fragment). Let us consider for simplicity the approach of a hydrogen atom to a proton. As the separation is decreased exchange takes place and the classical concept is no longer applicable. This happens just where the covalent energy becomes appreciable, and it is also at this point that the polarizability, which depends on virtual 1s + 2p transitions, decreases rapidly because of the promotion of the 2p orbitals to higher states in the united atom (see reference 1, particularly Fig. 1, p. 676). It is important to know how the polarization energy varies with r over the whole range as this determines how and where the potential energy curve is affected. This in turn determines the nature and position of spectra involving highly excited states of the neutral molecule which approach those of the molecular ion. In the case of Hz+ the term #ae2/r-4 equals 4.8 eV for r = 1 A, and as the polarization energy for this molecule is calculated by quantum theory to be O-46 eV it is clear that the effect of polarization is not fully operating around the equilibrium separation of l-06 A. In fact as r goes from infinity to zero the term w/t’ becomes equal to 046 eV
327
W. C. PRICE, P. V. HARRIS and T. R. PASSMORE
328
at 1.8 A. Thus we might expect the polarization energy to increase as l/r4 as the hydrogen atom is brought near the proton until the separation is just less than 2 A after which it will remain fairly constant due to the promotion of the p orbitals already mentioned. Fig. 1 shows the curves for H,+ . The polarization reduces the ionization energy of the Hi
H;:
3-
2z* g
2CV
I
2 H-H,
CH’ :
3
a
‘fl
I .o
C-H,
8
Fro. 1. E&et of polarization on potential eoerlly eurvcs. (u) polarization energy, (b) uncorrected, (c) corrected for polarization.
molecule by O-46 eV from what it would otherwise be.. Similar reductions are obtained when a non-bonding electron is removed from a hydride HX, the values of I(X)-I(HX) being 0*54,0~43,0~62,0~27,0~22,0~10 for F, 0, C, Cl, Br, I respectively. The flatness of the polarization energy curve means that it does not effect the vibration frequency or the equilibrium internuclear distance which are mainly controlled by covalent forces, i.e. it does not distort the potential energy curve in the neighbourhood of r. but just causes it to go on climbing at fairly large distances. It can and does affect the anharmonic constants of the higher vibrational states and causes them to change suddenly as in the case of the III state of CH+c2) (Fig. 1) as a result of the nature of the change in the polarization contribution to the potential energy curve. The unique character of this spectrum permits observation of the part of the potential energy curve where the covalent contribution becomes fractional and the polarization energy becomes dominant. Anomalous curves are also found in Nl and CO+. (Note: the 1lI curve of Fig. 1 has not been corrected in accordance with the new ionization potential of CH.)
The ionization and dissociation energies of molecules and radicals
329
When instead of a hydrogen atom a methyl group is attached to X+ it is to be expected that because of the higher polarizability of this group the contribution of polarization energy to the binding of the molecular ion would be correspondingly increased. This will be associated with a larger reduction in the ionization energy of the neutral molecule than that for its H analogue relative to the ionization energy of X-alternatively thought of as a larger increase in the bond dissociation energy of the molecular ion. A factor partly offsetting this could be the greater nuclear separation, though because of the flatness of the polarization energy curve about re this factor is not expected to have a large effect (effects depending on the electronegativity of X are also present, i.e. mixing of some XH+ with X+H which necessitate using a weighted mean of aH and ax). As a typical example the data for methyl chloride and HCl are given below in terms of a square array of values: Dissociation energy of CHs-X : Dissociation Ionization energy of CHs-X : Ionization 443 : 4.70 For HCl we have and for methyl chloride 12.74 : 13.01
energy of CHs-X+ . energy of X * 3 -40 : 5.13 -. 11.28 : 13.01
The four figures are of course related by the fact that the sums of diagonally opposite quantities are equal. It can be seen that the polarization binding of the methyl group to X+ in CHsX+ leads to an increase of 1.73 eV for the CX bond as compared with an increase of only 0.27 eV in the corresponding HX process. This is a factor of over six whereas the ratio of the polarizability of C to H is only 2.4 : 1 and the distance factor would tend to affect the ratio in the opposite sense.* It would be better to restrict ourselves to atoms in the first row of the periodic table and to consider only relatively nonpolar molecules in order to avoid the effect on a of contributions from ionic structures. We can use Ha0 and (CH&C whose I.P. loweiings relative to OsP are 1-O (13*61- 12.61) and 3.61 (13*61- 10.00) eV respectively, i.e. a ratio of 3*6/l, or again comparing CHs and the tertiary butyl radical relative to C?P, we have 1.42 (11.26 - 984) and 3.83 (11.26 7.43) eV, a ratio of 2*7/l (references 3 and 4). The appearance of the absorption bands of the higher Rydberg states of methyl chloride which is an example of the removal of a non-bgnding electron, show by the absence of vibrational structure that the CC1 bond distance and frequency are little altered by ionization t5). The increase in C-Cl dissociation energy of over 50 per cent all takes place at larger internuclear distances and does not affect the shape of the potential energy curve around re as is to be expected from the nature of the space variation of the polarization energy (cf. Nsf, CO+). Exactly the same explanation applies to the general features of the Rydberg bands resulting from the removal of a bonding electron, for example, a rr electron from ethylene@). Because of the polarization binding in the ion there is little difference between D(HsC = CHs) and D(HsC = CH2)+ the values of which are 5.40 and 5.29 eV respectively-the difference depends on I(CzHd)-I(CH2) = IO.51 - 1040 = 0.11 eV. Nevertheless the absorption spectra of the Rydberg states going to the ion show that the C = C frequency has been reduced from 1623 to 1370 cm-l and that there has been an t Other minor factors are D bond polarization, hybridization and the polarizability of X. In the heavier atoms X of lower ionization potential ground state structures of the type X+H- become important, cf. Has, HsSe, HsTe, and cause the ionization energies of HnX to be greater than that 6f X.
330
W. C. PRICE,P. V. HARRISand T. R. PASSMORE
increase in the internuclear distance as a result of excitation and ionization. For a neutral molecule this would be associated with a lowering of the C- C dissociation energy. However, because the polarization energy comes in at large values or r, the potential energy curve keeps rising until it reaches a dissociation energy little different from that of the ground state of the neutral molecule in spite of the fact that its form around the minimum is very different in the neutral and ionized molecule. When an antibonding electron is removed which is the case in 0s going to the lowest state of Os+‘ then because of the polarization energy the increase in bond energy in going to the ion is greater than might’ otherwise have been (6*48- 5.08 = 1.4 eV). Allowing for spin stabilization in 0s (1.1 eV) and polarizationin Os+(l eV) the antibonding power of the lnlrelectron can be estimated to be 1.5 eV. In the case of fluorine for which the polarizability is much lower, the compensating polarization energy is correspondingly less and a smaller increase (2.7 - 1.6 = 1.1 eV) is observed. No abnormal shape is evident in the potential curve of Or +. The reason for this is that the remaining contribution of the remaining antibonding electron varies in much the same way at large r as the polarization energy. Because of their opposite sign these cancel each other out over this range. Systematic study of the electronic structures of molecules In order to study systematically the electronic structures of molecules and to use the large amount of accurate spectroscopic, photo-ionization and thermochemical data which is now available, it is convenient as a Grst step to confine ourselves to atoms with only s and p electrons, i.e. to atoms in the first row of the periodic table, and to consider their combination with (a) H atoms and (b) with F atoms. The following Table 1 gives the dissociation energies of the ground states of such molecules. TABLE1. DLSOCMTION HYDRIDES
AND
(6)
ELEMENTS
OF
PERIODIC
TABLE.
THE
ENBRGIFS FLUORIDES
FIRST (%UUFS
ROW IN
OF (a) OF
THE
OF
THE
ev).
(a) Li, 2.2; Be, 2.5; B, 3.39; C, 341; N, 3.8; 0, 4.35; F, 5-87 (b) Li, 6.0; Be, 5.4; B, 8-5 ; C, 4.6 ; N, 3 ; 0, 2
; F, 1.63
It can be seen that the values for the hydrides rise gradually to HF particularly as more p electrons accumulate and make the bond more polar. From lithium to boron fluoride we have the gradual withdrawal of the s orbital from participation in bond formation (through s, sp to pure p) until in BF we have ~271.4 closed shells of bonding orbitals in which the w bond is formed from 2p orbitals. This gives the most stable structure and as further electrons have to go into antibonding G orbitals a steady reduction occurs in the dissociation energy until the low value of 1.63 eV is reached in Fs with four antibonding r electrons. In order to study the contribution to the structure which each electron makes it is desirable to remove it and to find what happens to the dissociation energy of the molecule making due allowance for any polarization energy or other effects which might become involved. For this purpose a table has been constructed of best values found in the literature and many from the authors’ laboratory. These are given as the square arrays of
331
The ionization and dissociation energies of molecules and radicals
dissociation and ionization energies which have been previously mentioned (Table 2). The values are of course related by the equation IRH+ D*R&
= IR+DRH
Thereisnot scope in this present article to discuss all the points in this table; a detailed discussion will be published elsewhere. A few of the salient features will however be pointed out. In nearly all cases D* is greater than DO due to the effect of polarization energy in the ion. In CF this difference is very large but this is attributable to the closed shell configuration of CF+. Also in CF2 the D* of FC - F+ is low because of the stable configuration of the product CF+. In CF4 +, D* is zero because the product CFs+ has a closed shell AND IONIZATION ENERGIES TABLE2. TABLEOF DL%oCL4TlON
-
I
4.476Ha 2648 15.423 13595
DQO~RH:D*~~RH+
Hea
I.P. of RH : I.P. of R (values in eV)
_-
--
0 21.5
3.1 24.58
BeH
LiH 2.3 8.4
(2) 5.39 LiF
3.2 9.32
c2.7) B;*8) (9.4)
--110.3)
_-
I
BeF
(6.2)
115.5
6’;“1
(3) 5.39
;jo
-
-
(85:94)
(
75) (13.8)
I
4.10 11.265
CHa 4.45 5.60 10.396 1064
7.4 11.265
(6.0) 11.1
($8;
4.5 10.15
5.7 11.3
CH 3.47 1064 CF
CHs 3-90 9.84
CH4
4.46 10,396
CF4
CFs
S:81
1.28 9.84
4.42 12.98
*
5.3 15.4
(3.9) cF:4.9) 11.1 10.1
1091
I
NH (3.7) (13.9)
OH
NHa (4.0) 11.3
(4.3) 14.545
(6.6) (13.9)
440 13.1
OHa 4.9 13.615
NFs (3.4) 11.9
(4.8) 14.545 -
5.6 13.1
(1.8) 13.6
OFs (0.7) (12.5)
FH 5.87 (6.5) (16.8) 17.42
-
__
NF (3.1) (12.8)
5.10 1260
-
(4.3) (12.8)
NFs (222 $98’
(2.2) 0F(3.3) 13,615 (12.5)
F8 1.63 16.3
2.7 17.42
-_
Most of the values given in this table have been taken from (1) G. HI~RzRRRG, Spectra of Diutomic Molecules, Van Nostrand, New York (1951), (2) A. G. GAYDON,Dissociation Energies, Chapman-Hall, London (1953), (3) T. L. CJTTIELL, The Strengths of Chemical Bon&, Butterworths, London (1958). D(BH) is from A. C,HURLEY, Proc. Roy. Sot. A261,237 (1961). The I.P. of CH has been kindly supplied by Dr. Her&erg (see also G. HERZBERG,Cur&. J. Phys., 39, 155 (1961).). The I.P. of CP is from J. W. C. JOHNSand R. F. BARROW,Froc. Phys. Sot., A71,478 (1958). that of CFs from J. L. MARGRAVE, J. Chem. Phys., 31,1432 (1959). Data on NF,, is from J. T. HERRONand V. H. DIBELER,J. Cherri.Phys., 35, 747 (1961) also reference 7. and on I.P. (Fs) from R. P. ICZKOWS~Iand J. L. ~%ARCJRAVE, J. Chem. Phys., 30, 403 (1959)-data reinterpreted. Other 1.P.s are from unpublished work in this laboratory. Values given in brackets are estimated by a method of isoelectronic similarity, details of which are to be published. Some I.P. values are from K. WATANABE,J. Chem. Phys., 26,542 (1957) whose measurements we have con8rmed independently.
W. C. PRKE,P. V. HARRISand T. R. PASSMORE
332
configuration in Dsr, symmetry. These last two facts account for the enormous difference between the fragmentation in mass spectrometry of saturated hydrocarbons and that of fluorocarbons. In the former the parent ions are present in appreciable abundance whereas in the latter they are never found being unstable and dissociating into the stable closed shell species CFf and CFsf. The situation is entirely different in the unsaturated fluorocarbons and arises largely from the relative stability of the CF2 radical. This stability is acquired by the orientation of the 2p C orbital along the Csv axis so avoiding antibonding interaction with T F electrons. Spectroscopic and photoionization work on these fluorinated hydrocarbons confirm all the features discussed above.
Bond deformation
energy
The very large deformation energies associated with polyatomic radicals having 71electrons is of great importance in chemistry. Where these form ‘bonding closed shell’ coplaner structures, great stability is achieved and this gives rise to the acidity and basicity of the ions, e.g. carbonate, buanidinium. In these molecular ions two n electrons go into ~1, the most strongly bonding of the (nr%ra4)set of orbitals available to the planar structure. There is great stability as this orbital spreads over the whole of the molecule having only one nodal plane which passes through all the nuclei. Any bending out of the plane of the structures would give nodal planes through bonds and convert this strongly bonding orbital to an antibonding or at least a non-bonding one. In CFsf where the 7~2 orbitalsi are also filled, the forces making for planarity are so large that in the formation of CF4+ by bringing up F to CFs+ more energy is required to deform the CFs+ into the required pyramidal configuration than would be released by the covalent energy available in a single electron FsC+ - F bond. This accounts for instability of CF4+. Similar explanations apply to the instability of CaFs+. Even in CFsH, where CFsH+ is only just stable, the CH bond is reduced to such an extent (O-6 eV) as compared with its value in CH4+ (l-28 eV) that this ion only appears in mass spectrometry with low abundance. From the above it can readily be appreciated that energies involved in distorting closed shell planar structures of 7~electrons are of the order one or more electron volts. Similar values are obtained thermochemically for the resonance energies of planar polyatomic anions and cations. This figure is also confirmed from studies of the spectra of polyatomic molecules and radicals in which one of the states is planar and the other non-planar. Such spectra usually extend over ca. 1 eV from the long wavelength limit to the absorption maximum. As the latter corresponds by the Franck-Condon principle to the geometrical configuration of the initial state and the former to that of the final state, their difference is equal to the energy required to alter the geometry of the final to that of the initial state. In contrast the deformation energies of analogous molecules with H instead of F substitution are only about a few tenths of a volt as is evident from the potential barrier of NHs which is 0.25 eV. The fact that CHs is planar means that it must have a barrier less than the zero point energy of the out-of-plane bending mode (i.e. N + (1000) cm-l or N 0.06 eV). The indications are that there is only a small difference in the energies t np has one additional nodal plane which is perpendicular to the molecular plane.
The ionization and dissociation energies of molecules and radicals
and this also supports hypothesis of CH deformation A theoretical value of 0.4 eV is given by JORDANand LONGUET-HIGGINS(~)and we deduce values of 0.32 and 0.16 eV for CHs and CHs+ from the ionization potentials of the radicals. The deformation energy in CHs+ is much greater than in CHs where the additional p electron weakens resistance to distortion. Its removal probably strengthens the planar structure so that about 1.8 eV would be required to distort it to a tetrahedral form. For CHs the value is O-4 eV. We are grateful to Dr. Herzberg for informing us of the ionization potential of 10.64 eV for CH. It is believed that all the figures in the CH, section of Table 2 are now accurate to a few hundredths of an electron volt. (They imply also that D(HsC! = CHs) = 5.40 eV or 127 kcals.) We are pleased to acknowledge support from the Department of Scientific and Industrial Research, Imperial Chemical Industries and the U.S. Department of the Army through its European Research Office under Contract DA-91-591 EUC-1683. between energies
CHs linear and 1.41 be&
333
of this low magnitude.
REFERENCES 1. J. E. LENNARD-JONES.Trans. Farad. Sot. 25,668 (1929). 2. A. E. DOUGLAS and J. R. MORTON, Astrophys. J. 131, 1 (1960). 3. G. HERZBERG,Proc. Roy. Sot. A 262, 291 (1961). 4. F. P. Lossx~o and J. B. DESOUSA,J. Amer. Chem. Sot. 81,281 (1959). 5. W. C. PRICE,J. Chem. Phys. 4,539 (1936). 6. W. C. PRICEand W. T. Tvrq Proc. Roy. Sot. A 174,207 (1940). 7. P. C. H. JORDANand H. C. UNCXIET-HIGGINS, Molec. Phys. 5, 121 (1962).
l'==L%~td~
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THE IONIZATION AND DISSOCIATION ENERGIES OF MOLECULES AND RADICALS W. C. PRICE, P. V. HARRIS and T. R. PASSMORE King’s College, London. W.C.2, England Abstract-It is shown that in molecular ions two factors, (i) polarization and (ii) radical deiormation energies, assume as large a role as covalent and ionic binding do in neutral mok.culea. This is illustrated particularly with reference to species XI& and XFn, where X refers to atoms in the Grst row of the periodic table. Values are given for the ionization and dissociation energies of all these molecules or radicals and their ions. The unusual space dependence of polarization energy is reconciled with the
appearance of Rydberg spectra, the vibration frequencies, and interatomic distances in the ionized state. IN THE ground states of neutral molecules the main factors responsible for molecular formation are the covalent and ionic binding between the atoms. When a molecule is dissociated work has to be done against these forces, the nature of which and their variation with internuclear separation is fairly well understood. The covalent forces which arise from the overlap of bonding orbitals fall off very rapidly, whereas the ionic forces are coulombic and decrease more slowly with distance, a factor which considerably affects the shapes of the corresponding potential energy curves. In the case of the ions of molecules or radicals two additional major binding factors are involved which become of the same order as those already mentioned as of major importance in neutral molecules. These are (a) polarization energy and (b) bond deformation energy. The latter is also of importance for neutral molecules. A molecular ion is formed from a charged and a neutral radical and as these approach, the polarization of the neutral molecule gives rise to a polarization energy which is equal to &~e2/r”’at fairly large distances (a being the polarizability of the neutral fragment). Let us consider for simplicity the approach of a hydrogen atom to a proton. As the separation is decreased exchange takes place and the classical concept is no longer applicable. This happens just where the covalent energy becomes appreciable, and it is also at this point that the polarizability, which depends on virtual 1s + 2p transitions, decreases rapidly because of the promotion of the 2p orbitals to higher states in the united atom (see reference 1, particularly Fig. 1, p. 676). It is important to know how the polarization energy varies with r over the whole range as this determines how and where the potential energy curve is affected. This in turn determines the nature and position of spectra involving highly excited states of the neutral molecule which approach those of the molecular ion. In the case of Hz+ the term #ae2/r-4 equals 4.8 eV for r = 1 A, and as the polarization energy for this molecule is calculated by quantum theory to be O-46 eV it is clear that the effect of polarization is not fully operating around the equilibrium separation of l-06 A. In fact as r goes from infinity to zero the term w/t’ becomes equal to 046 eV
327
W. C. PRICE, P. V. HARRIS and T. R. PASSMORE
328
at 1.8 A. Thus we might expect the polarization energy to increase as l/r4 as the hydrogen atom is brought near the proton until the separation is just less than 2 A after which it will remain fairly constant due to the promotion of the p orbitals already mentioned. Fig. 1 shows the curves for H,+ . The polarization reduces the ionization energy of the Hi
H;:
3-
2z* g
2CV
I
2 H-H,
CH’ :
3
a
‘fl
I .o
C-H,
8
Fro. 1. E&et of polarization on potential eoerlly eurvcs. (u) polarization energy, (b) uncorrected, (c) corrected for polarization.
molecule by O-46 eV from what it would otherwise be.. Similar reductions are obtained when a non-bonding electron is removed from a hydride HX, the values of I(X)-I(HX) being 0*54,0~43,0~62,0~27,0~22,0~10 for F, 0, C, Cl, Br, I respectively. The flatness of the polarization energy curve means that it does not effect the vibration frequency or the equilibrium internuclear distance which are mainly controlled by covalent forces, i.e. it does not distort the potential energy curve in the neighbourhood of r. but just causes it to go on climbing at fairly large distances. It can and does affect the anharmonic constants of the higher vibrational states and causes them to change suddenly as in the case of the III state of CH+c2) (Fig. 1) as a result of the nature of the change in the polarization contribution to the potential energy curve. The unique character of this spectrum permits observation of the part of the potential energy curve where the covalent contribution becomes fractional and the polarization energy becomes dominant. Anomalous curves are also found in Nl and CO+. (Note: the 1lI curve of Fig. 1 has not been corrected in accordance with the new ionization potential of CH.)
The ionization and dissociation energies of molecules and radicals
329
When instead of a hydrogen atom a methyl group is attached to X+ it is to be expected that because of the higher polarizability of this group the contribution of polarization energy to the binding of the molecular ion would be correspondingly increased. This will be associated with a larger reduction in the ionization energy of the neutral molecule than that for its H analogue relative to the ionization energy of X-alternatively thought of as a larger increase in the bond dissociation energy of the molecular ion. A factor partly offsetting this could be the greater nuclear separation, though because of the flatness of the polarization energy curve about re this factor is not expected to have a large effect (effects depending on the electronegativity of X are also present, i.e. mixing of some XH+ with X+H which necessitate using a weighted mean of aH and ax). As a typical example the data for methyl chloride and HCl are given below in terms of a square array of values: Dissociation energy of CHs-X : Dissociation Ionization energy of CHs-X : Ionization 443 : 4.70 For HCl we have and for methyl chloride 12.74 : 13.01
energy of CHs-X+ . energy of X * 3 -40 : 5.13 -. 11.28 : 13.01
The four figures are of course related by the fact that the sums of diagonally opposite quantities are equal. It can be seen that the polarization binding of the methyl group to X+ in CHsX+ leads to an increase of 1.73 eV for the CX bond as compared with an increase of only 0.27 eV in the corresponding HX process. This is a factor of over six whereas the ratio of the polarizability of C to H is only 2.4 : 1 and the distance factor would tend to affect the ratio in the opposite sense.* It would be better to restrict ourselves to atoms in the first row of the periodic table and to consider only relatively nonpolar molecules in order to avoid the effect on a of contributions from ionic structures. We can use Ha0 and (CH&C whose I.P. loweiings relative to OsP are 1-O (13*61- 12.61) and 3.61 (13*61- 10.00) eV respectively, i.e. a ratio of 3*6/l, or again comparing CHs and the tertiary butyl radical relative to C?P, we have 1.42 (11.26 - 984) and 3.83 (11.26 7.43) eV, a ratio of 2*7/l (references 3 and 4). The appearance of the absorption bands of the higher Rydberg states of methyl chloride which is an example of the removal of a non-bgnding electron, show by the absence of vibrational structure that the CC1 bond distance and frequency are little altered by ionization t5). The increase in C-Cl dissociation energy of over 50 per cent all takes place at larger internuclear distances and does not affect the shape of the potential energy curve around re as is to be expected from the nature of the space variation of the polarization energy (cf. Nsf, CO+). Exactly the same explanation applies to the general features of the Rydberg bands resulting from the removal of a bonding electron, for example, a rr electron from ethylene@). Because of the polarization binding in the ion there is little difference between D(HsC = CHs) and D(HsC = CH2)+ the values of which are 5.40 and 5.29 eV respectively-the difference depends on I(CzHd)-I(CH2) = IO.51 - 1040 = 0.11 eV. Nevertheless the absorption spectra of the Rydberg states going to the ion show that the C = C frequency has been reduced from 1623 to 1370 cm-l and that there has been an t Other minor factors are D bond polarization, hybridization and the polarizability of X. In the heavier atoms X of lower ionization potential ground state structures of the type X+H- become important, cf. Has, HsSe, HsTe, and cause the ionization energies of HnX to be greater than that 6f X.
330
W. C. PRICE,P. V. HARRISand T. R. PASSMORE
increase in the internuclear distance as a result of excitation and ionization. For a neutral molecule this would be associated with a lowering of the C- C dissociation energy. However, because the polarization energy comes in at large values or r, the potential energy curve keeps rising until it reaches a dissociation energy little different from that of the ground state of the neutral molecule in spite of the fact that its form around the minimum is very different in the neutral and ionized molecule. When an antibonding electron is removed which is the case in 0s going to the lowest state of Os+‘ then because of the polarization energy the increase in bond energy in going to the ion is greater than might’ otherwise have been (6*48- 5.08 = 1.4 eV). Allowing for spin stabilization in 0s (1.1 eV) and polarizationin Os+(l eV) the antibonding power of the lnlrelectron can be estimated to be 1.5 eV. In the case of fluorine for which the polarizability is much lower, the compensating polarization energy is correspondingly less and a smaller increase (2.7 - 1.6 = 1.1 eV) is observed. No abnormal shape is evident in the potential curve of Or +. The reason for this is that the remaining contribution of the remaining antibonding electron varies in much the same way at large r as the polarization energy. Because of their opposite sign these cancel each other out over this range. Systematic study of the electronic structures of molecules In order to study systematically the electronic structures of molecules and to use the large amount of accurate spectroscopic, photo-ionization and thermochemical data which is now available, it is convenient as a Grst step to confine ourselves to atoms with only s and p electrons, i.e. to atoms in the first row of the periodic table, and to consider their combination with (a) H atoms and (b) with F atoms. The following Table 1 gives the dissociation energies of the ground states of such molecules. TABLE1. DLSOCMTION HYDRIDES
AND
(6)
ELEMENTS
OF
PERIODIC
TABLE.
THE
ENBRGIFS FLUORIDES
FIRST (%UUFS
ROW IN
OF (a) OF
THE
OF
THE
ev).
(a) Li, 2.2; Be, 2.5; B, 3.39; C, 341; N, 3.8; 0, 4.35; F, 5-87 (b) Li, 6.0; Be, 5.4; B, 8-5 ; C, 4.6 ; N, 3 ; 0, 2
; F, 1.63
It can be seen that the values for the hydrides rise gradually to HF particularly as more p electrons accumulate and make the bond more polar. From lithium to boron fluoride we have the gradual withdrawal of the s orbital from participation in bond formation (through s, sp to pure p) until in BF we have ~271.4 closed shells of bonding orbitals in which the w bond is formed from 2p orbitals. This gives the most stable structure and as further electrons have to go into antibonding G orbitals a steady reduction occurs in the dissociation energy until the low value of 1.63 eV is reached in Fs with four antibonding r electrons. In order to study the contribution to the structure which each electron makes it is desirable to remove it and to find what happens to the dissociation energy of the molecule making due allowance for any polarization energy or other effects which might become involved. For this purpose a table has been constructed of best values found in the literature and many from the authors’ laboratory. These are given as the square arrays of
331
The ionization and dissociation energies of molecules and radicals
dissociation and ionization energies which have been previously mentioned (Table 2). The values are of course related by the equation IRH+ D*R&
= IR+DRH
Thereisnot scope in this present article to discuss all the points in this table; a detailed discussion will be published elsewhere. A few of the salient features will however be pointed out. In nearly all cases D* is greater than DO due to the effect of polarization energy in the ion. In CF this difference is very large but this is attributable to the closed shell configuration of CF+. Also in CF2 the D* of FC - F+ is low because of the stable configuration of the product CF+. In CF4 +, D* is zero because the product CFs+ has a closed shell AND IONIZATION ENERGIES TABLE2. TABLEOF DL%oCL4TlON
-
I
4.476Ha 2648 15.423 13595
DQO~RH:D*~~RH+
Hea
I.P. of RH : I.P. of R (values in eV)
_-
--
0 21.5
3.1 24.58
BeH
LiH 2.3 8.4
(2) 5.39 LiF
3.2 9.32
c2.7) B;*8) (9.4)
--110.3)
_-
I
BeF
(6.2)
115.5
6’;“1
(3) 5.39
;jo
-
-
(85:94)
(
75) (13.8)
I
4.10 11.265
CHa 4.45 5.60 10.396 1064
7.4 11.265
(6.0) 11.1
($8;
4.5 10.15
5.7 11.3
CH 3.47 1064 CF
CHs 3-90 9.84
CH4
4.46 10,396
CF4
CFs
S:81
1.28 9.84
4.42 12.98
*
5.3 15.4
(3.9) cF:4.9) 11.1 10.1
1091
I
NH (3.7) (13.9)
OH
NHa (4.0) 11.3
(4.3) 14.545
(6.6) (13.9)
440 13.1
OHa 4.9 13.615
NFs (3.4) 11.9
(4.8) 14.545 -
5.6 13.1
(1.8) 13.6
OFs (0.7) (12.5)
FH 5.87 (6.5) (16.8) 17.42
-
__
NF (3.1) (12.8)
5.10 1260
-
(4.3) (12.8)
NFs (222 $98’
(2.2) 0F(3.3) 13,615 (12.5)
F8 1.63 16.3
2.7 17.42
-_
Most of the values given in this table have been taken from (1) G. HI~RzRRRG, Spectra of Diutomic Molecules, Van Nostrand, New York (1951), (2) A. G. GAYDON,Dissociation Energies, Chapman-Hall, London (1953), (3) T. L. CJTTIELL, The Strengths of Chemical Bon&, Butterworths, London (1958). D(BH) is from A. C,HURLEY, Proc. Roy. Sot. A261,237 (1961). The I.P. of CH has been kindly supplied by Dr. Her&erg (see also G. HERZBERG,Cur&. J. Phys., 39, 155 (1961).). The I.P. of CP is from J. W. C. JOHNSand R. F. BARROW,Froc. Phys. Sot., A71,478 (1958). that of CFs from J. L. MARGRAVE, J. Chem. Phys., 31,1432 (1959). Data on NF,, is from J. T. HERRONand V. H. DIBELER,J. Cherri.Phys., 35, 747 (1961) also reference 7. and on I.P. (Fs) from R. P. ICZKOWS~Iand J. L. ~%ARCJRAVE, J. Chem. Phys., 30, 403 (1959)-data reinterpreted. Other 1.P.s are from unpublished work in this laboratory. Values given in brackets are estimated by a method of isoelectronic similarity, details of which are to be published. Some I.P. values are from K. WATANABE,J. Chem. Phys., 26,542 (1957) whose measurements we have con8rmed independently.
W. C. PRKE,P. V. HARRISand T. R. PASSMORE
332
configuration in Dsr, symmetry. These last two facts account for the enormous difference between the fragmentation in mass spectrometry of saturated hydrocarbons and that of fluorocarbons. In the former the parent ions are present in appreciable abundance whereas in the latter they are never found being unstable and dissociating into the stable closed shell species CFf and CFsf. The situation is entirely different in the unsaturated fluorocarbons and arises largely from the relative stability of the CF2 radical. This stability is acquired by the orientation of the 2p C orbital along the Csv axis so avoiding antibonding interaction with T F electrons. Spectroscopic and photoionization work on these fluorinated hydrocarbons confirm all the features discussed above.
Bond deformation
energy
The very large deformation energies associated with polyatomic radicals having 71electrons is of great importance in chemistry. Where these form ‘bonding closed shell’ coplaner structures, great stability is achieved and this gives rise to the acidity and basicity of the ions, e.g. carbonate, buanidinium. In these molecular ions two n electrons go into ~1, the most strongly bonding of the (nr%ra4)set of orbitals available to the planar structure. There is great stability as this orbital spreads over the whole of the molecule having only one nodal plane which passes through all the nuclei. Any bending out of the plane of the structures would give nodal planes through bonds and convert this strongly bonding orbital to an antibonding or at least a non-bonding one. In CFsf where the 7~2 orbitalsi are also filled, the forces making for planarity are so large that in the formation of CF4+ by bringing up F to CFs+ more energy is required to deform the CFs+ into the required pyramidal configuration than would be released by the covalent energy available in a single electron FsC+ - F bond. This accounts for instability of CF4+. Similar explanations apply to the instability of CaFs+. Even in CFsH, where CFsH+ is only just stable, the CH bond is reduced to such an extent (O-6 eV) as compared with its value in CH4+ (l-28 eV) that this ion only appears in mass spectrometry with low abundance. From the above it can readily be appreciated that energies involved in distorting closed shell planar structures of 7~electrons are of the order one or more electron volts. Similar values are obtained thermochemically for the resonance energies of planar polyatomic anions and cations. This figure is also confirmed from studies of the spectra of polyatomic molecules and radicals in which one of the states is planar and the other non-planar. Such spectra usually extend over ca. 1 eV from the long wavelength limit to the absorption maximum. As the latter corresponds by the Franck-Condon principle to the geometrical configuration of the initial state and the former to that of the final state, their difference is equal to the energy required to alter the geometry of the final to that of the initial state. In contrast the deformation energies of analogous molecules with H instead of F substitution are only about a few tenths of a volt as is evident from the potential barrier of NHs which is 0.25 eV. The fact that CHs is planar means that it must have a barrier less than the zero point energy of the out-of-plane bending mode (i.e. N + (1000) cm-l or N 0.06 eV). The indications are that there is only a small difference in the energies t np has one additional nodal plane which is perpendicular to the molecular plane.
The ionization and dissociation energies of molecules and radicals
and this also supports hypothesis of CH deformation A theoretical value of 0.4 eV is given by JORDANand LONGUET-HIGGINS(~)and we deduce values of 0.32 and 0.16 eV for CHs and CHs+ from the ionization potentials of the radicals. The deformation energy in CHs+ is much greater than in CHs where the additional p electron weakens resistance to distortion. Its removal probably strengthens the planar structure so that about 1.8 eV would be required to distort it to a tetrahedral form. For CHs the value is O-4 eV. We are grateful to Dr. Herzberg for informing us of the ionization potential of 10.64 eV for CH. It is believed that all the figures in the CH, section of Table 2 are now accurate to a few hundredths of an electron volt. (They imply also that D(HsC! = CHs) = 5.40 eV or 127 kcals.) We are pleased to acknowledge support from the Department of Scientific and Industrial Research, Imperial Chemical Industries and the U.S. Department of the Army through its European Research Office under Contract DA-91-591 EUC-1683. between energies
CHs linear and 1.41 be&
333
of this low magnitude.
REFERENCES 1. J. E. LENNARD-JONES.Trans. Farad. Sot. 25,668 (1929). 2. A. E. DOUGLAS and J. R. MORTON, Astrophys. J. 131, 1 (1960). 3. G. HERZBERG,Proc. Roy. Sot. A 262, 291 (1961). 4. F. P. Lossx~o and J. B. DESOUSA,J. Amer. Chem. Sot. 81,281 (1959). 5. W. C. PRICE,J. Chem. Phys. 4,539 (1936). 6. W. C. PRICEand W. T. Tvrq Proc. Roy. Sot. A 174,207 (1940). 7. P. C. H. JORDANand H. C. UNCXIET-HIGGINS, Molec. Phys. 5, 121 (1962).