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NITROGEN DIOXIDE AND DINITROGEN TETROXIDE
340
NITROGEN : K. JONES 12.5. A N A L Y S I S
Dinitrogen trioxide can be detected and determined even in the presence of its dissociation products NO and NO2, by infrared and ultraviolet spectroscopic methods510. Frequently, however, only an estimate of the total nitrogen oxide content is required, and methods based on their absorption as nitrite or nitrate followed by colorimetric analysis are employed531. Such methods are commonly used for analysis of "nitrous fumes"532 which usually refers to mixtures of NO and NO2 and other oxides that exist in equilibrium with them (principally N2O3 and N 2 0 4 ). However, the full analysis of such mixtures often requires a combination of techniques including physical equipment such as gas chromatographs using helium carrier gas and thermal conductivity detectors, or mass spectrometers531.
13. NITROGEN DIOXIDE AND DINITROGEN TETROXIDE 13.1. PRODUCTION AND INDUSTRIAL USE
Nitrogen dioxide and dinitrogen tetroxide are discussed together because under ordinary conditions of temperature and pressure both species occur in the presence of each other in a state of equilibrium533-5 : 2N0 2 ^N 2 0 4 Thus N2O4 having the higher molecular weight is the low-temperature form, and NO2 the dissociation product at higher temperatures. Their equilibrium concentrations at various temperatures are indicated in Table 71. Like other molecules with an odd number of electrons (free radicals), the monomer NO2 is coloured, whereas the dimeric form N2O4 is colourless, hence the intensity of the red-brown colour in the gas phase can be used to monitor the monomer concentration535. TABLE 71. EQUILIBRIUM CONCENTRATIONS OF NO2 AND N2O4 AT ORDINARY TEMPERATURES
Temperature
CO
Phase
-11.2 -11.2 21.15 21.15 135
Solid Liquid Liquid Gas Gas
531 K . I. V a s u , in ref. 459, p . 633.
532
Equilibrium concentration NO2
N204
0 0.01 0.1 15.9 99
100 99.99 99.9 84.1 1
Methods for the Detection of Toxic Substances in Air, Booklet No. 5, Nitrous Fumes, Ministry of Employment and Productivity, HM Factory Inspectorate, HMSO, London (1969). 533 C. C. Addison, Chemistry in Liquid Dinitrogen Tetroxide, Vol. Ill, Part I, Chemistry in Nonaqueous Ionizing Solvents (eds. G. Jander, H. Spandau and C. C. Addison), Pergamon, London (1967). 534 I. R. Beattie, in ref. 459, Section XXVII, pp. 246-68. « s p. Gray and A. D. Yoffe, Chem. Rev. 55 (1955) 1069.
PREPARATION OF NO2/N2O4
341
On an industrial scale, NO2-N2O4 is produced by the oxidation of nitric oxide during the manufacture of nitric acid 53 *: 2NO(g) + 0 2 (g)-2N0 2 (g) Thus the principal use of NO2-N2O4 in forming nitric acid 536 » 537 is described by its reaction with water according to the overall equation 3N02(g) + H20(1) ^ 2HN03(aq) + NO(g) 13.2. PREPARATION OF N 0 2 / N 2 0 4 The equilibrium mixture NO2-N2O4 is usually prepared on a laboratory scale either by reduction or decomposition of nitric acid, or more conveniently by the thermal decomposition of a heavy metal nitrate. Rigorously dried lead nitrate can be heated in a steel bomb 533 or a tube furnace 538 under a slow stream of oxygen which helps prevent contamination by N2O3: 2Pb(N03)2 -► 4N02+2PbO + 0 2 The evolved gases are passed through a condenser to remove traces of nitric acids then over phosphorus pentoxide to complete the drying process. The gas is condensed at — 78 °C and can be further purified by fractional distillation, rejecting the first fraction and collecting the remainder in individual ampoules which are then sealed. Pure dinitrogen tetroxide freezes as a colourless solid, whereas a pale-green solid results if traces of moisture are present, colour being a sensitive test of purity 533 . Other methods which have been used to prepare nitrogen dioxide include the oxidation of nitric oxide: 2NO + 0 2 - * 2 N 0 2 the reaction of dinitrogen trioxide and dinitrogen pentoxide: N203 + N205-*4N02 the reaction of nitric acid with sulphur dioxide: 2HN03 + S0 2 -* 2N0 2 +H 2 S0 4 the reaction of nitric acid and phosphorus pentoxide: 4HN0 3 + 2P205 -> 4N0 2 +0 2 +4HP0 3 the reaction of nitric acid and copper: Cu+4HN0 3 -> 2N0 2 +Cu(N0 3 ) 2 +2H 2 0 the reaction of nitrosyl chloride with silver nitrate: NOC1+AgN03 -> 2N02+AgCl the reaction of nitrosonium hydrogen sulphate and potassium nitrate: NOHS0 4 +KN0 3 -> 2N0 2 +KHS0 4 536 1. R. Beattie, in ref. 459, Section XXII, pp. 158-72. C. Matasa and E. Matasa, L'Industrie Moderne des Produits Azotes, Dunod, Paris (1968), p. 460. 53 8 P. W. Schenk, in Handbook of Preparative Inorganic Chemistry, 2nd edn., Vol. I (ed. G. Brauer), Academic Press, New York (1963), p. 488. 537
342
NITROGEN: K. JONES
Such methods may be convenient and adequate for some purposes, but the presence of other non-metal compounds generally leads to impurities in the product. 13.3. NUCLEAR ISOTOPES OF NITROGEN DIOXIDE*** The oxidation of enriched nitric oxide with excess molecular oxygen is rapid and complete at room temperature : i5NO+i02->15N02
Similarly, N^O^O and N18C>2 can be prepared by treating N^O with normal and 1 8 0 enriched oxygen respectively. Based on the thermal decomposition of heavy metal nitrates, 15NC>2 has been prepared by heating K15NC>3 with PbC>2, or dissolving K15NC>3 in 85% phosphoric acid and adding copper metal: M(i5N0 3 ) 2 -> 4 i 5 N 0 2 + 2 M O + 0 2
Nitrogen dioxide enriched in 1 8 0 can also be prepared540 by heating the lead nitrate that is formed from lead chloride and KN18C>3. The nitrogen oxides collected at — 196°C are vaporized and allowed to react with oxygen to ensure complete oxidation of the nitric oxide to the nitrogen dioxide. NO2 is then re-condensed at — 78 °C, and the excess oxygen and nitrogen pumped off. 13.4. STRUCTURE AND BONDING The solid phase consists entirely of N2O4 molecules and shows two modifications known as the monoclinic formai which is unstable with respect to the cubic form54*. The structures of the molecules in both phases which have been determined by X-ray diffraction are similar and are in agreement with a planar structure with a centre of symmetry: O
O
\
/
N—N \ / O O although infrared evidence suggests that non-planar O2N-NO2 and even ONONO2 structures are present at — 196°C and — 269°C respectively54^. Indeed, nitrosonium nitrate NO+NO 3 has been prepared by oxidation of NO at — 196°C and it has been suggested544 that the stabilities of the N2O4 isomers decrease in the order 02NN02 > ΝΟ+ΝΟ3 > ONON02. Structural evidence for N2O4 in the liquid phase is less direct, but it is thought that more than one form exists and in particular one through which the formation of ion pairs [NO+][NO]7 can be accounted533. Electron diffraction study of the vapour at — 20°C 5
39 W . Spindel, in Inorganic Isotopic Syntheses (ed. R . H . H e r b e r ) , Benjamin, N e w Y o r k (1962), p . 101. I. Dostrovsky and D. Samuel, in Inorganic Isotopic Syntheses (ed. R. H. Herber), Benjamin, New York (1962), p. 134. 541 P . G r o t h , Acta chem. scand. 17 (1963) 2419. 542 j . s . Broadley a n d J. M . R o b e r t s o n , Nature 164 (1949) 915. 543 I. c . Hisatsune a n d J. P . Devlin, / . Chem. Phys. 3 1 (1959) 1130. 544 L . Parts a n d J. Y . Miller, J r . , / . Chem. Phys. 4 3 (1965) 136. 540
343
STRUCTURE AND BONDING OF N 0 2
supports the planar structure5^ although there seems to be no obvious reason why there should not be free rotation of the NO2 groups about the long N-N bond. The V-shaped structure of nitrogen dioxide in the gas phase has been confirmed by a variety of spectroscopic methods and by electron diffraction*^ giving the N-O bond length of 1.19 Â and the ONO bond angle of 134°. Using Linnett-type valence bond structures, electronic configurations which maintain a complete octet of electrons round each atom are possible for the seventeen-electron NO2 molecule547 as shown in Fig. 31.
. ^v.
.0.
.0,
—
. y^.
..q:
.(?.
FIG. 31. Valence bond structures of NO2.
The bonding can also be described in terms of molecular orbital theory, and N 0 2 provides a good example to illustrate the procedure by which molecular orbitals can be constructed for simple molecules548. Since NO2 is angular, the orbitals of the molecule must be characterized according to the operations of the point group which describes the molecular shape Civ for which the character table is:
C2v
E
Ciiz)
σν(χζ)
Ai A2 Bi B2
1 1 1 1
1 1 -1 -1
1 -1 1 -1
1 -1 -1 1
The NO2 molecule is fixed in the co-ordinate system so as to be situated say in the xz plane as illustrated in Fig. 32. By carrying out the operations of the point group C2v[Ci(z\ συ(χζ) etc.] on each atomic orbital of nitrogen (2s, 2pz, etc.) and then on linear combinations of the orbitals on the oxygen atoms [(.?+£), (s-s), (pzApz), etc.], the transformation properties of each orbital can be calculated and assigned to an irreducible representation (au a2i etc., corresponding 545 R. G. Snyder and I. C. Hisatsune, / . Chem. Phys. 26 (1957) 960. 546 G. R. Bird, J. Chem. Phys. 25 (1956) 1040.
547
M. Green, in Developments in Inorganic Nitrogen Chemistry, Vol. I (ed. C. B. Colburn), Elsevier, Amsterdam (1966), pp. 1-71. 548 C. J. Ballhausen and H. B. Gray, Molecular
Orbital Theory, Benjamin, N e w York (1965), p. 76.
344
NITROGEN: K. JONES
FIG. 32. The N 0 2 molecule.
to Ai, A2 in the character table). Thus having assigned the available atomic orbitals to the appropriate symmetry type, the molecular orbitals for NO2 can be constructed simply by combining those atomic orbitals with the same symmetry. These molecular orbitals can then be presented in the form of an energy-level diagram as in Fig. 33 into which electrons E
i
2p .hi \\
d2
AA \
2s
*i> \
lb
\\ w
2//
^-v^·
(Pz-Pz) (Ρχ-Ρχ) (Py-Py) 2p (Py + Py) (Px-Px) (Pz+Pz) J
2b, y;
Y
\
w -
2a,
\\
\ ^ ^ b
N atomic orbitals
N02 molecular orbitals
.
J
(s-s)!
2s
(s + s)]
-*
U
O O atomic orbital combinations
FIG. 33. Molecular orbital diagram for NO2.
STRUCTURE AND BONDING OF N 0 2
345
are allocated to the lowest molecular orbitals to give the ground state configuration of the molecule: *AX (1αι)2(161)2(2Λι)2(2^1)2(1^2)2(3Λι)2(1ΰ2)2(361)2(4αι)ΐ
Thus although the 1#2, 36ι and 4αι levels are all fairly close together, the unpaired electron is assigned to the 4a\ orbital if one considers the large splitting that is observed in the e.s.r. spectrum. This is produced by an unpaired s electron interacting with the spin of the 14N nucleus, and both la 2 and 3b\ are made up of/? functions, whereas the 4a\ level is composed partly of nitrogen and oxygen 2s orbitals. The first excited state has the configuration: 22?! (lfli)2(161)2(2a1)2(261)2(162)2(3a1)2(lfl2)2(361)i(4ai)2
03 a 117° r 1.28
NO. 134°~ 1.19
C02 180° 1.16
FIG. 34. Ozone, nitrogen dioxide and carbon dioxide correlation diagrams.
and the absorption band in the N 0 2 spectrum at 23,000 cm -1 is usually assigned to the Mi -► 2#i transition547,548. A diagram similar to that in Fig. 33 can be arrived at by correlating the molecular orbital diagrams of structurally related compounds. From this point of view, NO2 can be considered as the compromise situation between ozone, a C2,, molecule with smaller angle and longer bonds, and carbon dioxide, a linear molecule with point group Dœh and shorter bonds. The correlation diagram in Fig. 34 indicates the approximate energies of NO2 molecular orbitals simply by joining the appropriate corresponding orbitals of O3 and CO2 (note that the different notation for CO2 is due to it having a centre of symmetry). These
346
NITROGEN : K. JONES
correlation diagrams can also be used to estimate the bond angles of low-lying excited states while the bond length can be inferred by considering to what extent antibonding orbitals are involved549. With regard to the description of the bonding in N2O4, this would appear to be less straightforward533. There is general agreement that the molecule is planar, has a long N-N bond and all the N-O bonds are of equal length. Planarity in this situation implies a restriction to rotation about the N-N bond which normally would be expected to result from steric restraint or multiple bonding. However, there would not appear to be any steric hindrance, particularly in the gas phase, and a bond length longer than that of a normal N-N single bond (N-N in hydrazine is 1.47 Â) is not consistent with multiple bonding either. Thus this unusual situation has led to the proposal of various novel descriptions of the bonding. The usual structure is a resonance hybrid of the canonical forms illustrated in Fig. 35.
°\
o·
+ -N .+/
7°
\
Ov —
—
V.
J
P
N — &' N
X
X - K FIG. 35. Valence bond structures for N2O4.
N ·N
/ "V
:0
t
-*
*-
O:
/ \ N · N
/
N · N
:θ
-*
»-
v
Q.
./ v N · N
FIG. 36. Linnett non-pairing structures for N2O4. 549 A. D. Walsh, J. Chem. Soc. (1953) 2266.
STRUCTURE AND BONDING OF N 2 0 4
347
Here the long N-N bond is explained on the basis of repulsion due to formal positive charges lying on adjacent atoms, but it does not offer a satisfactory explanation for planarity. Similarly, resonance between the Linnett non-pairing structures in Fig. 36 also provides a simple picture of N 2 0 4 which nicely explains the long one-electron N-N bond but fails to account for the planarity547. More recently it was suggested that the central bond is a π bond with little or no σ character550. However, this is not now thought to be correct, and calculations551 suggest TABLE 72. MOLECULAR CONSTANTS OF N2O4
Property X-ray diffraction dataa System N - N bond length N-O bond length ONO bond angle Electron diffraction datab on gaseous N2O4 N - N bond length N-O bond length ONO bond angle Fundamental vibration frequencies'" for liquid N2O4
v, 1
n > (a,) "3 J V4
M
Z }<*·>
V7 V8
(blu) (b2g)
::„ }<**>
:; K>
Enthalpy of homolytic dissociation Enthalpy of heterolytic dissociation Entropy of homolytic dissociation Entropy of heterolytic dissociation Dipole moment Dielectric constant Magnetic susceptibility (293) Ultraviolet absorption maxima Refractive index (5893 Â) a b
Ref. 542. Ref. 545.
Value
Monoclinic (unstable) 1.75 A 1.21 A 135°
Cubic (stable) 1.64 Â 1.17 A 126°
1.75 A 1.18 A 133.7° (cm-i) 1379.6 808 260 (-50) 1712 482 429 622 1748 381 1262 750 12.9 kcal/mol 7.87 kcal/mol 42 cal/deg mol -80.1 cal/deg mol 0.55 D 2.42 2.52 x 10~7 c.g.s. units g _1 343 τημ 1.420 c d
Ref. 552. Ref. 533.
550 c . A. Coulson and J. Duchesne, Bull. Acad. Belg. CL Sei. 43 (1957) 522. 551 M. Green and J. W. Linnett, Trans. Faraday Soc. 57 (1961) 10.
348
NITROGEN I K. JONES
that although the main part of the N-N bond is σ in character, there is a fairly large π contribution from orbitals which lie in the plane of the molecule (similar to combinations of the 3b2 orbitals of NO2) and not from the more usual orbitals perpendicular to the plane. 13.5. MOLECULAR CONSTANTS Some molecular constants for N2O4 are listed in Table 72, while values for NO2 and NO 2 are compared in Table 73 together with some relevant spectroscopic552 and thermodynamic data553. Electron spin resonance spectra have been recorded for NO2 in various matrices554. T A B L E 73. M O L E C U L A R C O N S T A N T S 3 O F NO2
Structural p a r a m e t e r s N - O b o n d length O N O b o n d angle F u n d a m e n t a l vibration frequencies b Vl ( û l ) V2 ( « l )
Magnetic susceptibility ESR hyperfine splitting (NO2 trapped in argon) Ionization potential Dipole moment Thermodynamic functions0 Δ / / / 0 (enthalpy of formation (0 K) Δ///0 (enthalpy of formation) (298 K) #298-#o (enthalpy) A G/0 (Gibbs energy of formation) (298 K) S° (entropy) Cp (heat capacity) a b c
NOj
N02
Property
V3 ( £ l )
AND
1.197 Â 134° 15'
NOJ
1.15 A 180°
1322.5 cm-i 749.7 cm-i 1617.75 cm-i 28.2 x 10-2 c.g.s. u n its g-i 58 oersted 11 eV 0.39/) 8.6 kcal/mol 7.93 kcal/mol 2.438 kcal/mol
276.6 kcal/mol 277.4 kcal/mol
12.26 kcal/mol 57.35 cal/deg mol 8.89 cal/deg mol
Refs. 533-5. Ref. 552. Ref. 553.
13.6. P H Y S I C A L P R O P E R T I E S
Some physical data for N 0 2 have already been mentioned in Table 73, but so few experimentally determined values are known because of their need to be measured at temperatures above 200°C to ensure complete dissociation of N 2 0 4 . Values of physical constants for N 2 0 4 are listed in Table 74. 552 R. F . B a r r o w a n d A. J. Merer, in ref. 459, p . 515. 553 D . D . W a g m a n , W . H . E v a n s , V. B . P a r k e r , I. H a l o w , S. M . Bailey a n d R . H . S c h u m m , Selected Values of Chemical Thermodynamic Properties, N B S Technical N o t e 270-3, Washington (1968). 554 p . w . Atkins a n d M . C. R . Symons, The Structure of Inorganic Radicals, Elsevier, A m s t e r d a m (1967), p. 127.
PHYSICAL PROPERTIES OF DINITROGEN TETROXIDE
349
However, only the solid phase consists of pure N2O4, and values for properties of the liquid refer to material containing the equilibrium quantity of NO2. This probably makes no difference just above the melting point but is likely to be significant towards the boiling point, accompanying the colour change from very pale yellow to deep red-brown over the same temperature range. The solvent properties of liquid N2O4 will be discussed further in section 13.7.2. TABLE 74. PHYSICAL PROPERTIES OF DINITROGEN TETROXIDE a Property
Freezing point Boiling point Critical temperature Critical pressure Critical density Critical volume Δ # fusion Entropy of fusion AH vaporization Entropy of vaporization Specific heat Specific heat Cryoscopic constant Ebullioscopic constant Thermodynamic properties0 AH/Q (enthalpy of formation) (g) ΔΗ/0 (enthalpy of formation) (g) H29B-H0 (enthalpy) (g) ΔΗ/ο (enthalpy of formation) (1) Δ Gf (Gibbs energy of formation) (g) Δ Gf (Gibbs energy of formation) (1) So (entropy) (g) So (entropy) (1) CPo (heat capacity) (g) Cpo (heat capacity) (1) Vapour pressure of solid Vapour pressure of liquid
Density of solid Density of liquid
Temperature (K) 261.95 294.30 158.2
261.95 261.95 294.30 294.30 261.95 294.30
0 298
Value
-11.2°C 21.15°C -114.95°C 99.96 atm 0.570 g/cm3 165.3 cnvVmol 3.502 kcal/mol 13.37 cal/deg mol 9.110 kcal/mol 30.96 cal/deg mol 32.65 cal/deg mol 33.90 cal/deg mol 3.64°C depression/kg N 2 0 4 1.37°C elevation/kg N 2 0 4 4.49 kcal/mol 2.19 kcal/mol 3.918 kcal/mol — 4.66 kcal/mol 23.38 kcal/mol 23.29 kcal/mol 72.70 cal/deg mol 50.0 cal/deg mol 18.47 cal/deg mol 34.1 cal/deg mol
298 298 298 298 298 298 298 240.3-261.9 l o g i 0 j p = 9.58149- 2460.000/Γ+7.61700/Γ3-1.51335Γ2/10 5 261.9-294.9 l o g i o P = 8.00436- 1753.000/Γ-11.8078/104+2.0954Γ2/106 248 70 mm 261.95 139.8 mm 266 mm 272.31 454 mm 283 569.4 mm 288 294.30 760 mm 1.979 g/cm3 78 253-294.30 d= 1.4927-2.235//103-2.75/2/106 253 1.5364 g/cm3 258 1.5256 g/cm3 1.5147 g/cm3 263 1.4927 g/cm3 273 1.4702 g/cm3 283 1.4469 g/cm3 293
350
NITROGEN: K. JONES TABLE 74 (cont.)
Property
Viscosity of liquid
Electrical conductivity of liquid
a
Ref. 533.
b
Value
Temperature (K) 253-294.3 253 258 263 273 283 293 248-294.3 248 290
logioi/ = 400/T-1.742 0.687 poise 0.641 poise 0.599 poise 0.527 poise 0.468 poise 0.420 poise l o g / c = - 1 2 6 7 / Γ - 8.260 4.3 x 10-14 Ω - i c m - i 2.36 x 10-13 Ω - i c m - i
Ref. 553.
13.7. C H E M I S T R Y A N D C H E M I C A L P R O P E R T I E S
13.7.1. Reactions of N 0 2 It is unlikely that perfectly pure NO2 can be formed at atmospheric pressure, for the temperature required to dissociate N2O4 also coincides with the onset of thermal decomposition of NO2 which becomes significant above 150°C and complete536 at about 600°C: N2O4 ^ 2N02 ^ 2NO + 0 2 At low partial pressures, photochemical decomposition occurs and can be studied under conditions where the recombination of nitric oxide and oxygen is so slow as to be negligible. The proposed mechanism555 is: N02+ÄV->NOJ
NO? + N02->2NO + 0 2 Low concentrations of NO2 in hydrogen and oxygen-air mixtures depress considerably the ignition temperatures and effect the explosion limits556. The overall reaction between NO2 and hydrogen is given as N02 + H 2 -^NO + H20 for which the following mechanism has been proposed557: H2 + N0 2 ->H+HN0 2 H+N0 2 ->NO + OH OH+H 2 ^H 2 0 + H OH+N02 + M -> HNO3 + M OH + NO + M->HN02 + M 555 R . G . W . N o r r i s h , / . Chem. Soc. (1929) 1158. 556 A. C. E g e r t o n a n d J. Powling, Proc. Roy. Soc. A , 193 (1948) 172. 557 p . G . A s h m o r e a n d B. P . Levitt, Trans. Faraday Soc. 53 (1957) 945.
CHEMISTRY AND CHEMICAL PROPERTIES OF N 0 2
351
Such reactions are also thought to be involved in the reactions between water and NO2 in the production of nitric acid536: 2 N 0 2 + H 2 0 -► HNO3+HNO2 3 H N 0 2 -> HNO3 + 2NO + H2O
It is not surprising therefore that N 0 2 in the presence of those compounds with which it is in equilibrium is strongly corrosive towards metals. Other oxidations involving nitrogen dioxide558 include the reactions with halogen compounds : NO2 + 2HCI - * NOCl + H 2 0 + i C l 2
At higher temperatures the nitrosyl halide decomposes: N02 + 2HC1 -> NO + H20 + Cl2 N0 2 + 2HBr -> NO + H20 + Br2 Likewise, a variety of bimolecular gas phase reactions involving NO2 have been studied55^: NO2+NH3 -> HNO2+NH2 N02 + F2-^N0 2 F+F NO2 + CI2O -> NO2CI + OCI N0 2 + C102->N03 + 0C1 N0 2 + CO->NO + C02 Thus the description of nitrogen dioxide as a fairly reactive free radical is well illustrated by dimerization, combination with other free radicals560, addition, and abstraction reactions, with inorganic and organic molecules561. 13.7.2. Reactions of Liquid Dinitrogen Tetroxide533 Homolytic dissociation of N2O4 always occurs to a small extent in the liquid phase: N 2 0 4 (1) ^ 2N0 2 (1);
Δ/Zdiss = 19.5 kcal/mol
but of the various alternative heterolytic dissociation processes, it is perhaps surprising that the nitronium and nitrite ions have never been recognized as free ions in liquid N 2 0 4 . In fact the chemistry of the N2O4 solvent system is rationalized on the basis of self-ionization to give nitrosyl and nitrate ions according to the equation N 2 0 4 ^NO+ + N03 analogous to the ionization of water or liquid ammonia: 2H 2 O^H 3 0 + + OH-
558 j . H . T h o m a s , in Oxidation and Combustion Reviews, Vol. I (ed. C . F . H . Tipper), Elsevier, Amsterd a m (1965), p . 137. 559 A . F . T r o t m a n - D i c k e n s o n a n d G . S. Milne, Tables of Bimolecular Gas Reactions, N B S - 9 , W a s h i n g t o n (1969). 560 y . Rees a n d G . H . Williams, in Advances in Free-radical Chemistry, Vol. 3 (éd. G . H . Williams), L o g o s Press, L o n d o n (1969), p . 199. 561 A . V. Topchiev (translated by C. Matthews), Nitration of Hydrocarbons and other Organic Compounds, P e r g a m o n , L o n d o n (1959).
352
NITROGEN : K. JONES
Thus although the free ions in liquid N2O4 cannot be detected by methods such as infrared or Raman spectroscopy due to their low concentration, the large difference between the molar polarization and molar refraction is attributed562 to the presence of polar constituents in the liquid phase which are not present in either solid or gaseous states. These polar constituents are likely to be ion pairs rather than free ions because of the low dielectric constant of the solvent, and although ionic reactions will still occur, they do so as a result of collisions between ion pairs rather than free ions. Consequently, simple metal salts are insoluble in pure liquid N 2 0 4 , but do dissolve in mixtures of N2O4 diluted with solvents of higher dielectric constant533. Thus the addition of nitromethane (ε = 37) produces conducting solutions in which the self-ionization of N2O4 is enhanced to such an extent that dissociation is complete in solvents like perchloric or sulphuric acids: N2O4 + 3H2S04 -> NO+HS04 + HNO3 + SO3 + H30+HS04 Likewise a similar situation can be achieved by the addition of a donor solvent (such as ethyl acetate, diethyl ether, diethylnitrosamine, etc.). This arises from the electron-deficient nature of N2O4 and its tendency to form molecular addition compounds with donor molecules533, in particular oxygen and nitrogen bases bonded to the nitrogen(s) of N2O4: nD + N204 ^ D„ · N2O4 ^ (D„NO)+NO 3The wide variety of reactions in which liquid N2O4 can participate are summarized briefly below, although it should not be assumed that all such reactions are necessarily unique to the liquid phase. Acid-base reactions are described by the neutralization equation NO+ + N O i ^ N 2 0 4 Thus in liquid N2O4, nitrosyl chloride reacts with silver nitrate 563 : NOC1 + AgN0 3
N2o4
> AgCl + N 2 0 4
comparable with the corresponding reactions in water or liquid ammonia: HC1 + KOH H2°> KC1 + H20 NH4CI + NaNH2 —-i> NaCl + 2NH3 Similarly, compounds that produce "basic" solutions in liquid N2O4 are typified by alkylammonium nitrates which introduce excess nitrate into the system. Thus the reaction of zinc in these solutions illustrates the amphoteric behaviour of zinc compounds564 in N204: Ζη + 2ΕίΝΗ 3 ·Νθ3 + 2Ν 2 θ4 -* [EtNH 3 ] 2 [Zn(N0 3 )4] + 2NO
again directly analogous to reactions in water and liquid ammonia: Zn + 2NaOH + 2H 2 0 -> Na2[Zn(OH)4] + H 2 Zn + 2NaNH 2 +2NH 3 -> Na 2 [Zn(NH 2 )4]+H2 5
62 C. C. Addison, H. C. Bolton and J. Lewis, / . Chem. Soc. (1951) 1294. « C. C. Addison and R. Thompson, / . Chem. Soc. (1949) S211. 564 c . C. Addison and N. Hodge, / . Chem. Soc. (1954) 1138. 5
CHEMISTRY AND CHEMICAL PROPERTIES OF N 2 0 4
353
In contrast, the nitrosonium compounds are the "solvo-acids" and increase the concentrat i o n ^ of the NO + ion in liquid N2O4. Thus using tin as example, Ν2θ4
Sn+2NOC1
► SnCl2 + 2NO
again analogous to the corresponding reactions in water and liquid ammonia: Sn + 2HC1 - ^ - > SnCl2 + H 2 Sn + 2NH4C1
NH
% SnCl2 + 2NH 3 + H 2
Silver, zinc, mercury and the alkali metals react quite readily when brought into contact with liquid N2O4. For example, sodium requires continual stirring to expose clean metal surface if the reaction is to proceed to completion : Na + N204 -> NaN03 + NO At room temperature nitric oxide is evolved, but at lower temperatures or under pressure the formation of blue-green N2O3 is apparent533. The reaction is again directly analogous to those in water and liquid ammonia: 2Na+2H20 -> 2NaOH + H2 2Na + 2NH3 -» 2NaNH2 + H2 However, a number of other metals (Cd, Mn, Co, Cu, In, U) are only soluble if the liquid N2O4 is diluted with a solvent of higher dielectric constant such as nitromethane. The mixture becomes more reactive by enhancement of the self-ionization process of N2O4: MeN02
Cu + 3N204 > Cu(N03)2 · N204 + 2NO Similarly, the presence of donor solvents like ethyl acetate, diethyl ether, dimethyl sulphoxide, etc., in liquid N2O4 form the basis of excellent solvent systems for metals giving a variety of products including some stable solvate compounds containing dinitrogen tetroxide of crystallizations33: Ca+2N204 Mg+3N204 Bi+3N204 Cu + N 2 0 4
EtOAc EtOAc Me 2 SO AcOH 5-30% PhCN
> Ca(N0 3 ) 2 + 2NO > Mg(N03)2-N204+2NO > Bi(N0 3 ) 3 -3Me 2 SO + 3NO > Cu(OAc)2
Cu+5N204 > Cu(N03) · 2PhCN · 4N204 + NO Solvolysis reactions in liquid N2O4 are reactions of a salt MX with the solvent in which the anion X is totally or partially replaced by the nitrate ion, being the anion characteristic of the medium : MX + N2O4 -> MN03 + NOX Many such reactions have been reported including the following selection533: R4NC14- N 2 0 4 ^ R 4 N · N 0 3 + NOC1 ZnCl2 + N 2 0 4 ^ Zn(N0 3 ) 2 + 2NOC1 CaO + 2 N 2 0 4 -* Ca(N0 3 ) 2 + N 2 0 3 N a 2 0 2 + N 2 0 4 -► 2NaN0 3 N a O H + N 2 0 4 -> N a N 0 3 + H N 0 2
354
NITROGEN : K. JONES
This reaction provides an excellent route to anhydrous metal nitrates from the corresponding halides, particularly where X = Br or I as NOX decomposes preventing the possible formation of nitrosyl compounds533: Til4 + 4N204 ^ Ti(N03)4+4NO + 2I2 The oxidation properties of N2O4 towards inorganic compounds generally involve the donation of oxygen atoms rather than the removal of electrons from metal ions. This is well illustrated by the unusual salts that have been prepared by the oxidation of sodium hyponitrite565 with N 2 0 4 : Na2N202
rapid
> £-Na 2 N 2 0 3
slow
> Na 2 N 2 O s
slow
» Na2N206
100°
ΝΗ 2 ΟΗΊ EtN0 3 > -> a-Na 2 N 2 0 3
rapid
> Na2N204
slow
> Na 2 N 2 0 6
NaOEtJ Accordingly, the reaction of N2O4 with cobalt forms cobalt(II) rather than cobalt(III) compounds, and its selective oxidizing properties in organic chemistry566 are unique, particularly its ability to restrict oxidation to a single stage when further oxidation would appear to be possible. Dinitrogen tetroxide reacts readily with water according to the overall equation 3 N 2 0 4 + 2H 2 0 -> 4 H N 0 3 + 2NO
Several mechanisms, more than one of which may be applicable, have been proposed to account for the absorption stage in nitric acid production536 (see section 18.3) including the following: N 2 0 4 (g) -> N2O4O) N 2 0 4 (1) + H20(1) -> HN0 3 (1) + HN0 2 (1) 2HN02(1) -> HN0 3 (l) + NO(l) N 2 0 4 - > N O + + NOä N 2 0 4 + H 2 0 -> H + + N 0 J + H N 0 2 2 H N 0 2 -> H 2 0 + NO + N 0 2
However, in spite of this ready reaction, simple hydrated salts do not dehydrate completely when immersed in liquid N 2 0 4 (e.g. MgS0 4 -7H 2 0 -+ MgS0 4 -2H 2 0), and the selective reactivity may be related to possible differences in the bonding of the water567. Metal carbonyls react with N2O4 under mild conditions and provide simple preparative routes to metal nitrates and their N 2 0 4 solvates533: Co2(CO)8 + 8 N 2 0 4 -> 2Co(N0 3 ) 2 · 2 N 2 0 4 + 8CO + 4NO Fe(CO)5 + 4 N 2 0 4 -> Fe(N0 3 ) 3 · N 2 0 4 + 5CO + 3NO Mn2(CO)io + N 2 0 4 -> Mn(CO) 5 N0 3 + [Mn(CO)x(NO)y] 5
« C. C. Addison, G. A. Gamlen and R. Thompson, / . Chem. Soc. (1952) 338.
566 B. O. Field and J. Grundy, J. Chem. Soc. (1955) 1110. 567 C. C. Addison and D . J. Chapman, / . Chem. Soc. (1965) 819.
355
ANALYSIS OF N 0 2 - N 2 0 4
The reaction of N2O4 with organometallic compounds like tetramethyltin is explosively violent even at —80° and the reagents require dilution with an inert solvent5*^.· Me4Sn + 2 N 2 0 4
EtOAc > Me2Sn(N03)2
Finally, the reaction of N2O4 with the electron acceptor molecule boron trifluoride has received much attention. The reagents react both in 1:1 and 1:2 ratio to give stoichiometric solid products which have been formulated as containing the nitronium ion 567 : N2O4
BF3
> NOJ[BF3N02]-
BF3
> NOÎ[F3BONOBF3]-
Thus the overall scheme for N2O4 dissociation can be considered as in the following scheme533 : 2 N 0 2 ^ N 2 0 4 ^ [ N 0 2 + + NO7] ^ ™
» NO+ + NO3donor solvents
BF,
N O j [BF3-N02]"
13.8.
BIOLOGICAL
DnN204
(D n NO) + N0 3 -
PROPERTIES570
The toxic properties of nitrogen dioxide are well known, but the presence of dangerous concentrations in everyday situations is not widely appreciated. NO2 is the most dangerous component of "nitrous fumes", the name given to mixtures of oxides of nitrogen of variable composition. These mixtures can arise as a result of many industrial processes involving combustion but particularly electric-arc processes, electric-arc welding and oxy-acetylene welding. Hazards occur when such operations are carried out in confined, unventilated spaces. It is not considered safe to work for prolonged periods in an atmosphere containing >5ppm NO2, and short periods at higher concentrations can lead to symptoms such as inflammation of the lower respiratory tract and a marked elevation in the blood platelet count. Continued inhalation of concentrations > 100 ppm has led to death. 13.9.
ANALYSIS OF
N02-N204571
In large enough amounts, nitrogen dioxide is readily recognized by its characteristic red-brown colour. It can also be identified in equilibrium with its dimer by spectroscopic methods552; NO2 gives infrared absorption at 1626 cm -1 . Ultraviolet spectroscopy can also be used because dinitrogen tetroxide absorbs up to about 390 πιμ, and nitrogen dioxide in the 390-500 ιημ region. Nitrogen dioxide gives a positive test with Griess reagent due 568 c . C . A d d i s o n , W . B . Simpson a n d A . Walker, / . Chem. Soc. (1964) 2360. 569 R . w . Sprague, A . B . G a r r e t t a n d H . H . Sisler, / . Am. Chem. Soc. 82 (1960) 1059. 570 F. Call, in Mellows Comprehensive Treatise on Inorganic and Theoretical Chemistry, Vol. 8, Suppl. I, Nitrogen, P a r t I, L o n g m a n s , L o n d o n (1964), p p . 604-9. 571 K. I. Vasu, in Mellows Comprehensive Treatise on Inorganic and Theoretical Chemistry, Vol. 8, Suppl. I I , Nitrogen, P a r t I I , L o n g m a n s , L o n d o n (1967), p . 632. 572 M . J. T a r a s , in Colorimetric Determination of Nonmetals (ed. D . F . Boltz), Interscience, N e w Y o r k (1958).
356
NITROGEN: K. JONES
to the presence of nitrite in solution, which can also form the basis of a colorimetric method of analysis572. The application of titrimetric methods is somewhat similar to those described for nitric oxide. Nitrogen dioxide can be oxidized with excess cerium (IV) which is back-titrated with sodium oxalate: N 0 2 + Ce4+ + H 2 0 -> N O " + Ce3+ + 2 H +
Reference should also be made, however, to the analysis of nitrous fumes. This usually refers to mixtures of nitric oxide, nitrogen dioxide and other oxides that exist in equilibrium with them (principally N2O3 and N2O4). Their toxicity has directed considerable attention to the detection and estimation of small amounts of such materials573. Thus in this connection, and frequently in other situations, only the total nitrogen oxide content is required, and this can be obtained by oxidation with alkaline hydrogen peroxide to nitrate which can be determined colorimetrically by the phenoldisulphonic acid method. This can be preceded by absorption in Griess-Ilosvay reagent to obtain a separate estimation of N 0 2 content if required. For full analysis of such nitrogen oxide mixtures, mass spectrometry provides the quickest most reliable method. However, more recently, gas Chromatographie techniques have been developed which are capable of separating mixtures containing oxides of both nitrogen and carbon as well as elemental nitrogen using helium as carrier gas and thermal conductivity cell detectors574.
14. DINITROGEN PENTOXIDE 14.1. PREPARATION
Dinitrogen pentoxide is usually prepared in the laboratory by the dehydration of concentrated nitric acid with phosphorus pentoxide 575 · 576 : 2 H N O 3 + P 2 O 5 -> N2O5 + 2HPO3
Considerable care must be exercised during this reaction, as the procedure577 involves either the dropwise addition of concentrated HNO3 (d 1.525g/cm3) onto a large excess of P2O5 at — 10°C followed by slow distillation at about 35°C, or by adding P2O5 in one portion to nitric acid previously cooled to — 78 °C then allowing the mixture to attain room temperature slowly578. The principal alternative method is by the oxidation of nitrogen dioxide with ozone: 2 Ν 0 2 + θ 3 - * Ν 2 θ 5 + θ2 573 Methods for the Detection of Toxic Substances in Air, Booklet N o . 5, Nitrous Fumes, Ministry o f Employment and Productivity, H M Factory Inspectorate, H M S O , London (1969). 574 j . Janak and M. Rusek, Chemické Listy 48 (1954) 397. 5751. R. Beattie, in ref. 571, pp. 269-77. 576 p . w . Schenk, in Handbook of Preparative Inorganic Chemistry, Vol. I (ed. G. Brauer), Academic Press, N e w York (1963), p. 489. 577 N . S. Gruenhut, M. Goldfrank, M. L. Cushing and G. V. Caesar, Inorganic Syntheses, Vol. I l l (ed. L. F. Andrieth), McGraw-Hill, N e w York (1950), p. 78. ^78 G. V. Caesar and M. Goldfrank, / . Am. Chem. Soc. 68 (1962) 372.
NITROGEN : K. JONES 12.5. A N A L Y S I S
Dinitrogen trioxide can be detected and determined even in the presence of its dissociation products NO and NO2, by infrared and ultraviolet spectroscopic methods510. Frequently, however, only an estimate of the total nitrogen oxide content is required, and methods based on their absorption as nitrite or nitrate followed by colorimetric analysis are employed531. Such methods are commonly used for analysis of "nitrous fumes"532 which usually refers to mixtures of NO and NO2 and other oxides that exist in equilibrium with them (principally N2O3 and N 2 0 4 ). However, the full analysis of such mixtures often requires a combination of techniques including physical equipment such as gas chromatographs using helium carrier gas and thermal conductivity detectors, or mass spectrometers531.
13. NITROGEN DIOXIDE AND DINITROGEN TETROXIDE 13.1. PRODUCTION AND INDUSTRIAL USE
Nitrogen dioxide and dinitrogen tetroxide are discussed together because under ordinary conditions of temperature and pressure both species occur in the presence of each other in a state of equilibrium533-5 : 2N0 2 ^N 2 0 4 Thus N2O4 having the higher molecular weight is the low-temperature form, and NO2 the dissociation product at higher temperatures. Their equilibrium concentrations at various temperatures are indicated in Table 71. Like other molecules with an odd number of electrons (free radicals), the monomer NO2 is coloured, whereas the dimeric form N2O4 is colourless, hence the intensity of the red-brown colour in the gas phase can be used to monitor the monomer concentration535. TABLE 71. EQUILIBRIUM CONCENTRATIONS OF NO2 AND N2O4 AT ORDINARY TEMPERATURES
Temperature
CO
Phase
-11.2 -11.2 21.15 21.15 135
Solid Liquid Liquid Gas Gas
531 K . I. V a s u , in ref. 459, p . 633.
532
Equilibrium concentration NO2
N204
0 0.01 0.1 15.9 99
100 99.99 99.9 84.1 1
Methods for the Detection of Toxic Substances in Air, Booklet No. 5, Nitrous Fumes, Ministry of Employment and Productivity, HM Factory Inspectorate, HMSO, London (1969). 533 C. C. Addison, Chemistry in Liquid Dinitrogen Tetroxide, Vol. Ill, Part I, Chemistry in Nonaqueous Ionizing Solvents (eds. G. Jander, H. Spandau and C. C. Addison), Pergamon, London (1967). 534 I. R. Beattie, in ref. 459, Section XXVII, pp. 246-68. « s p. Gray and A. D. Yoffe, Chem. Rev. 55 (1955) 1069.
PREPARATION OF NO2/N2O4
341
On an industrial scale, NO2-N2O4 is produced by the oxidation of nitric oxide during the manufacture of nitric acid 53 *: 2NO(g) + 0 2 (g)-2N0 2 (g) Thus the principal use of NO2-N2O4 in forming nitric acid 536 » 537 is described by its reaction with water according to the overall equation 3N02(g) + H20(1) ^ 2HN03(aq) + NO(g) 13.2. PREPARATION OF N 0 2 / N 2 0 4 The equilibrium mixture NO2-N2O4 is usually prepared on a laboratory scale either by reduction or decomposition of nitric acid, or more conveniently by the thermal decomposition of a heavy metal nitrate. Rigorously dried lead nitrate can be heated in a steel bomb 533 or a tube furnace 538 under a slow stream of oxygen which helps prevent contamination by N2O3: 2Pb(N03)2 -► 4N02+2PbO + 0 2 The evolved gases are passed through a condenser to remove traces of nitric acids then over phosphorus pentoxide to complete the drying process. The gas is condensed at — 78 °C and can be further purified by fractional distillation, rejecting the first fraction and collecting the remainder in individual ampoules which are then sealed. Pure dinitrogen tetroxide freezes as a colourless solid, whereas a pale-green solid results if traces of moisture are present, colour being a sensitive test of purity 533 . Other methods which have been used to prepare nitrogen dioxide include the oxidation of nitric oxide: 2NO + 0 2 - * 2 N 0 2 the reaction of dinitrogen trioxide and dinitrogen pentoxide: N203 + N205-*4N02 the reaction of nitric acid with sulphur dioxide: 2HN03 + S0 2 -* 2N0 2 +H 2 S0 4 the reaction of nitric acid and phosphorus pentoxide: 4HN0 3 + 2P205 -> 4N0 2 +0 2 +4HP0 3 the reaction of nitric acid and copper: Cu+4HN0 3 -> 2N0 2 +Cu(N0 3 ) 2 +2H 2 0 the reaction of nitrosyl chloride with silver nitrate: NOC1+AgN03 -> 2N02+AgCl the reaction of nitrosonium hydrogen sulphate and potassium nitrate: NOHS0 4 +KN0 3 -> 2N0 2 +KHS0 4 536 1. R. Beattie, in ref. 459, Section XXII, pp. 158-72. C. Matasa and E. Matasa, L'Industrie Moderne des Produits Azotes, Dunod, Paris (1968), p. 460. 53 8 P. W. Schenk, in Handbook of Preparative Inorganic Chemistry, 2nd edn., Vol. I (ed. G. Brauer), Academic Press, New York (1963), p. 488. 537
342
NITROGEN: K. JONES
Such methods may be convenient and adequate for some purposes, but the presence of other non-metal compounds generally leads to impurities in the product. 13.3. NUCLEAR ISOTOPES OF NITROGEN DIOXIDE*** The oxidation of enriched nitric oxide with excess molecular oxygen is rapid and complete at room temperature : i5NO+i02->15N02
Similarly, N^O^O and N18C>2 can be prepared by treating N^O with normal and 1 8 0 enriched oxygen respectively. Based on the thermal decomposition of heavy metal nitrates, 15NC>2 has been prepared by heating K15NC>3 with PbC>2, or dissolving K15NC>3 in 85% phosphoric acid and adding copper metal: M(i5N0 3 ) 2 -> 4 i 5 N 0 2 + 2 M O + 0 2
Nitrogen dioxide enriched in 1 8 0 can also be prepared540 by heating the lead nitrate that is formed from lead chloride and KN18C>3. The nitrogen oxides collected at — 196°C are vaporized and allowed to react with oxygen to ensure complete oxidation of the nitric oxide to the nitrogen dioxide. NO2 is then re-condensed at — 78 °C, and the excess oxygen and nitrogen pumped off. 13.4. STRUCTURE AND BONDING The solid phase consists entirely of N2O4 molecules and shows two modifications known as the monoclinic formai which is unstable with respect to the cubic form54*. The structures of the molecules in both phases which have been determined by X-ray diffraction are similar and are in agreement with a planar structure with a centre of symmetry: O
O
\
/
N—N \ / O O although infrared evidence suggests that non-planar O2N-NO2 and even ONONO2 structures are present at — 196°C and — 269°C respectively54^. Indeed, nitrosonium nitrate NO+NO 3 has been prepared by oxidation of NO at — 196°C and it has been suggested544 that the stabilities of the N2O4 isomers decrease in the order 02NN02 > ΝΟ+ΝΟ3 > ONON02. Structural evidence for N2O4 in the liquid phase is less direct, but it is thought that more than one form exists and in particular one through which the formation of ion pairs [NO+][NO]7 can be accounted533. Electron diffraction study of the vapour at — 20°C 5
39 W . Spindel, in Inorganic Isotopic Syntheses (ed. R . H . H e r b e r ) , Benjamin, N e w Y o r k (1962), p . 101. I. Dostrovsky and D. Samuel, in Inorganic Isotopic Syntheses (ed. R. H. Herber), Benjamin, New York (1962), p. 134. 541 P . G r o t h , Acta chem. scand. 17 (1963) 2419. 542 j . s . Broadley a n d J. M . R o b e r t s o n , Nature 164 (1949) 915. 543 I. c . Hisatsune a n d J. P . Devlin, / . Chem. Phys. 3 1 (1959) 1130. 544 L . Parts a n d J. Y . Miller, J r . , / . Chem. Phys. 4 3 (1965) 136. 540
343
STRUCTURE AND BONDING OF N 0 2
supports the planar structure5^ although there seems to be no obvious reason why there should not be free rotation of the NO2 groups about the long N-N bond. The V-shaped structure of nitrogen dioxide in the gas phase has been confirmed by a variety of spectroscopic methods and by electron diffraction*^ giving the N-O bond length of 1.19 Â and the ONO bond angle of 134°. Using Linnett-type valence bond structures, electronic configurations which maintain a complete octet of electrons round each atom are possible for the seventeen-electron NO2 molecule547 as shown in Fig. 31.
. ^v.
.0.
.0,
—
. y^.
..q:
.(?.
FIG. 31. Valence bond structures of NO2.
The bonding can also be described in terms of molecular orbital theory, and N 0 2 provides a good example to illustrate the procedure by which molecular orbitals can be constructed for simple molecules548. Since NO2 is angular, the orbitals of the molecule must be characterized according to the operations of the point group which describes the molecular shape Civ for which the character table is:
C2v
E
Ciiz)
σν(χζ)
Ai A2 Bi B2
1 1 1 1
1 1 -1 -1
1 -1 1 -1
1 -1 -1 1
The NO2 molecule is fixed in the co-ordinate system so as to be situated say in the xz plane as illustrated in Fig. 32. By carrying out the operations of the point group C2v[Ci(z\ συ(χζ) etc.] on each atomic orbital of nitrogen (2s, 2pz, etc.) and then on linear combinations of the orbitals on the oxygen atoms [(.?+£), (s-s), (pzApz), etc.], the transformation properties of each orbital can be calculated and assigned to an irreducible representation (au a2i etc., corresponding 545 R. G. Snyder and I. C. Hisatsune, / . Chem. Phys. 26 (1957) 960. 546 G. R. Bird, J. Chem. Phys. 25 (1956) 1040.
547
M. Green, in Developments in Inorganic Nitrogen Chemistry, Vol. I (ed. C. B. Colburn), Elsevier, Amsterdam (1966), pp. 1-71. 548 C. J. Ballhausen and H. B. Gray, Molecular
Orbital Theory, Benjamin, N e w York (1965), p. 76.
344
NITROGEN: K. JONES
FIG. 32. The N 0 2 molecule.
to Ai, A2 in the character table). Thus having assigned the available atomic orbitals to the appropriate symmetry type, the molecular orbitals for NO2 can be constructed simply by combining those atomic orbitals with the same symmetry. These molecular orbitals can then be presented in the form of an energy-level diagram as in Fig. 33 into which electrons E
i
2p .hi \\
d2
AA \
2s
*i> \
lb
\\ w
2//
^-v^·
(Pz-Pz) (Ρχ-Ρχ) (Py-Py) 2p (Py + Py) (Px-Px) (Pz+Pz) J
2b, y;
Y
\
w -
2a,
\\
\ ^ ^ b
N atomic orbitals
N02 molecular orbitals
.
J
(s-s)!
2s
(s + s)]
-*
U
O O atomic orbital combinations
FIG. 33. Molecular orbital diagram for NO2.
STRUCTURE AND BONDING OF N 0 2
345
are allocated to the lowest molecular orbitals to give the ground state configuration of the molecule: *AX (1αι)2(161)2(2Λι)2(2^1)2(1^2)2(3Λι)2(1ΰ2)2(361)2(4αι)ΐ
Thus although the 1#2, 36ι and 4αι levels are all fairly close together, the unpaired electron is assigned to the 4a\ orbital if one considers the large splitting that is observed in the e.s.r. spectrum. This is produced by an unpaired s electron interacting with the spin of the 14N nucleus, and both la 2 and 3b\ are made up of/? functions, whereas the 4a\ level is composed partly of nitrogen and oxygen 2s orbitals. The first excited state has the configuration: 22?! (lfli)2(161)2(2a1)2(261)2(162)2(3a1)2(lfl2)2(361)i(4ai)2
03 a 117° r 1.28
NO. 134°~ 1.19
C02 180° 1.16
FIG. 34. Ozone, nitrogen dioxide and carbon dioxide correlation diagrams.
and the absorption band in the N 0 2 spectrum at 23,000 cm -1 is usually assigned to the Mi -► 2#i transition547,548. A diagram similar to that in Fig. 33 can be arrived at by correlating the molecular orbital diagrams of structurally related compounds. From this point of view, NO2 can be considered as the compromise situation between ozone, a C2,, molecule with smaller angle and longer bonds, and carbon dioxide, a linear molecule with point group Dœh and shorter bonds. The correlation diagram in Fig. 34 indicates the approximate energies of NO2 molecular orbitals simply by joining the appropriate corresponding orbitals of O3 and CO2 (note that the different notation for CO2 is due to it having a centre of symmetry). These
346
NITROGEN : K. JONES
correlation diagrams can also be used to estimate the bond angles of low-lying excited states while the bond length can be inferred by considering to what extent antibonding orbitals are involved549. With regard to the description of the bonding in N2O4, this would appear to be less straightforward533. There is general agreement that the molecule is planar, has a long N-N bond and all the N-O bonds are of equal length. Planarity in this situation implies a restriction to rotation about the N-N bond which normally would be expected to result from steric restraint or multiple bonding. However, there would not appear to be any steric hindrance, particularly in the gas phase, and a bond length longer than that of a normal N-N single bond (N-N in hydrazine is 1.47 Â) is not consistent with multiple bonding either. Thus this unusual situation has led to the proposal of various novel descriptions of the bonding. The usual structure is a resonance hybrid of the canonical forms illustrated in Fig. 35.
°\
o·
+ -N .+/
7°
\
Ov —
—
V.
J
P
N — &' N
X
X - K FIG. 35. Valence bond structures for N2O4.
N ·N
/ "V
:0
t
-*
*-
O:
/ \ N · N
/
N · N
:θ
-*
»-
v
Q.
./ v N · N
FIG. 36. Linnett non-pairing structures for N2O4. 549 A. D. Walsh, J. Chem. Soc. (1953) 2266.
STRUCTURE AND BONDING OF N 2 0 4
347
Here the long N-N bond is explained on the basis of repulsion due to formal positive charges lying on adjacent atoms, but it does not offer a satisfactory explanation for planarity. Similarly, resonance between the Linnett non-pairing structures in Fig. 36 also provides a simple picture of N 2 0 4 which nicely explains the long one-electron N-N bond but fails to account for the planarity547. More recently it was suggested that the central bond is a π bond with little or no σ character550. However, this is not now thought to be correct, and calculations551 suggest TABLE 72. MOLECULAR CONSTANTS OF N2O4
Property X-ray diffraction dataa System N - N bond length N-O bond length ONO bond angle Electron diffraction datab on gaseous N2O4 N - N bond length N-O bond length ONO bond angle Fundamental vibration frequencies'" for liquid N2O4
v, 1
n > (a,) "3 J V4
M
Z }<*·>
V7 V8
(blu) (b2g)
::„ }<**>
:; K>
Enthalpy of homolytic dissociation Enthalpy of heterolytic dissociation Entropy of homolytic dissociation Entropy of heterolytic dissociation Dipole moment Dielectric constant Magnetic susceptibility (293) Ultraviolet absorption maxima Refractive index (5893 Â) a b
Ref. 542. Ref. 545.
Value
Monoclinic (unstable) 1.75 A 1.21 A 135°
Cubic (stable) 1.64 Â 1.17 A 126°
1.75 A 1.18 A 133.7° (cm-i) 1379.6 808 260 (-50) 1712 482 429 622 1748 381 1262 750 12.9 kcal/mol 7.87 kcal/mol 42 cal/deg mol -80.1 cal/deg mol 0.55 D 2.42 2.52 x 10~7 c.g.s. units g _1 343 τημ 1.420 c d
Ref. 552. Ref. 533.
550 c . A. Coulson and J. Duchesne, Bull. Acad. Belg. CL Sei. 43 (1957) 522. 551 M. Green and J. W. Linnett, Trans. Faraday Soc. 57 (1961) 10.
348
NITROGEN I K. JONES
that although the main part of the N-N bond is σ in character, there is a fairly large π contribution from orbitals which lie in the plane of the molecule (similar to combinations of the 3b2 orbitals of NO2) and not from the more usual orbitals perpendicular to the plane. 13.5. MOLECULAR CONSTANTS Some molecular constants for N2O4 are listed in Table 72, while values for NO2 and NO 2 are compared in Table 73 together with some relevant spectroscopic552 and thermodynamic data553. Electron spin resonance spectra have been recorded for NO2 in various matrices554. T A B L E 73. M O L E C U L A R C O N S T A N T S 3 O F NO2
Structural p a r a m e t e r s N - O b o n d length O N O b o n d angle F u n d a m e n t a l vibration frequencies b Vl ( û l ) V2 ( « l )
Magnetic susceptibility ESR hyperfine splitting (NO2 trapped in argon) Ionization potential Dipole moment Thermodynamic functions0 Δ / / / 0 (enthalpy of formation (0 K) Δ///0 (enthalpy of formation) (298 K) #298-#o (enthalpy) A G/0 (Gibbs energy of formation) (298 K) S° (entropy) Cp (heat capacity) a b c
NOj
N02
Property
V3 ( £ l )
AND
1.197 Â 134° 15'
NOJ
1.15 A 180°
1322.5 cm-i 749.7 cm-i 1617.75 cm-i 28.2 x 10-2 c.g.s. u n its g-i 58 oersted 11 eV 0.39/) 8.6 kcal/mol 7.93 kcal/mol 2.438 kcal/mol
276.6 kcal/mol 277.4 kcal/mol
12.26 kcal/mol 57.35 cal/deg mol 8.89 cal/deg mol
Refs. 533-5. Ref. 552. Ref. 553.
13.6. P H Y S I C A L P R O P E R T I E S
Some physical data for N 0 2 have already been mentioned in Table 73, but so few experimentally determined values are known because of their need to be measured at temperatures above 200°C to ensure complete dissociation of N 2 0 4 . Values of physical constants for N 2 0 4 are listed in Table 74. 552 R. F . B a r r o w a n d A. J. Merer, in ref. 459, p . 515. 553 D . D . W a g m a n , W . H . E v a n s , V. B . P a r k e r , I. H a l o w , S. M . Bailey a n d R . H . S c h u m m , Selected Values of Chemical Thermodynamic Properties, N B S Technical N o t e 270-3, Washington (1968). 554 p . w . Atkins a n d M . C. R . Symons, The Structure of Inorganic Radicals, Elsevier, A m s t e r d a m (1967), p. 127.
PHYSICAL PROPERTIES OF DINITROGEN TETROXIDE
349
However, only the solid phase consists of pure N2O4, and values for properties of the liquid refer to material containing the equilibrium quantity of NO2. This probably makes no difference just above the melting point but is likely to be significant towards the boiling point, accompanying the colour change from very pale yellow to deep red-brown over the same temperature range. The solvent properties of liquid N2O4 will be discussed further in section 13.7.2. TABLE 74. PHYSICAL PROPERTIES OF DINITROGEN TETROXIDE a Property
Freezing point Boiling point Critical temperature Critical pressure Critical density Critical volume Δ # fusion Entropy of fusion AH vaporization Entropy of vaporization Specific heat Specific heat Cryoscopic constant Ebullioscopic constant Thermodynamic properties0 AH/Q (enthalpy of formation) (g) ΔΗ/0 (enthalpy of formation) (g) H29B-H0 (enthalpy) (g) ΔΗ/ο (enthalpy of formation) (1) Δ Gf (Gibbs energy of formation) (g) Δ Gf (Gibbs energy of formation) (1) So (entropy) (g) So (entropy) (1) CPo (heat capacity) (g) Cpo (heat capacity) (1) Vapour pressure of solid Vapour pressure of liquid
Density of solid Density of liquid
Temperature (K) 261.95 294.30 158.2
261.95 261.95 294.30 294.30 261.95 294.30
0 298
Value
-11.2°C 21.15°C -114.95°C 99.96 atm 0.570 g/cm3 165.3 cnvVmol 3.502 kcal/mol 13.37 cal/deg mol 9.110 kcal/mol 30.96 cal/deg mol 32.65 cal/deg mol 33.90 cal/deg mol 3.64°C depression/kg N 2 0 4 1.37°C elevation/kg N 2 0 4 4.49 kcal/mol 2.19 kcal/mol 3.918 kcal/mol — 4.66 kcal/mol 23.38 kcal/mol 23.29 kcal/mol 72.70 cal/deg mol 50.0 cal/deg mol 18.47 cal/deg mol 34.1 cal/deg mol
298 298 298 298 298 298 298 240.3-261.9 l o g i 0 j p = 9.58149- 2460.000/Γ+7.61700/Γ3-1.51335Γ2/10 5 261.9-294.9 l o g i o P = 8.00436- 1753.000/Γ-11.8078/104+2.0954Γ2/106 248 70 mm 261.95 139.8 mm 266 mm 272.31 454 mm 283 569.4 mm 288 294.30 760 mm 1.979 g/cm3 78 253-294.30 d= 1.4927-2.235//103-2.75/2/106 253 1.5364 g/cm3 258 1.5256 g/cm3 1.5147 g/cm3 263 1.4927 g/cm3 273 1.4702 g/cm3 283 1.4469 g/cm3 293
350
NITROGEN: K. JONES TABLE 74 (cont.)
Property
Viscosity of liquid
Electrical conductivity of liquid
a
Ref. 533.
b
Value
Temperature (K) 253-294.3 253 258 263 273 283 293 248-294.3 248 290
logioi/ = 400/T-1.742 0.687 poise 0.641 poise 0.599 poise 0.527 poise 0.468 poise 0.420 poise l o g / c = - 1 2 6 7 / Γ - 8.260 4.3 x 10-14 Ω - i c m - i 2.36 x 10-13 Ω - i c m - i
Ref. 553.
13.7. C H E M I S T R Y A N D C H E M I C A L P R O P E R T I E S
13.7.1. Reactions of N 0 2 It is unlikely that perfectly pure NO2 can be formed at atmospheric pressure, for the temperature required to dissociate N2O4 also coincides with the onset of thermal decomposition of NO2 which becomes significant above 150°C and complete536 at about 600°C: N2O4 ^ 2N02 ^ 2NO + 0 2 At low partial pressures, photochemical decomposition occurs and can be studied under conditions where the recombination of nitric oxide and oxygen is so slow as to be negligible. The proposed mechanism555 is: N02+ÄV->NOJ
NO? + N02->2NO + 0 2 Low concentrations of NO2 in hydrogen and oxygen-air mixtures depress considerably the ignition temperatures and effect the explosion limits556. The overall reaction between NO2 and hydrogen is given as N02 + H 2 -^NO + H20 for which the following mechanism has been proposed557: H2 + N0 2 ->H+HN0 2 H+N0 2 ->NO + OH OH+H 2 ^H 2 0 + H OH+N02 + M -> HNO3 + M OH + NO + M->HN02 + M 555 R . G . W . N o r r i s h , / . Chem. Soc. (1929) 1158. 556 A. C. E g e r t o n a n d J. Powling, Proc. Roy. Soc. A , 193 (1948) 172. 557 p . G . A s h m o r e a n d B. P . Levitt, Trans. Faraday Soc. 53 (1957) 945.
CHEMISTRY AND CHEMICAL PROPERTIES OF N 0 2
351
Such reactions are also thought to be involved in the reactions between water and NO2 in the production of nitric acid536: 2 N 0 2 + H 2 0 -► HNO3+HNO2 3 H N 0 2 -> HNO3 + 2NO + H2O
It is not surprising therefore that N 0 2 in the presence of those compounds with which it is in equilibrium is strongly corrosive towards metals. Other oxidations involving nitrogen dioxide558 include the reactions with halogen compounds : NO2 + 2HCI - * NOCl + H 2 0 + i C l 2
At higher temperatures the nitrosyl halide decomposes: N02 + 2HC1 -> NO + H20 + Cl2 N0 2 + 2HBr -> NO + H20 + Br2 Likewise, a variety of bimolecular gas phase reactions involving NO2 have been studied55^: NO2+NH3 -> HNO2+NH2 N02 + F2-^N0 2 F+F NO2 + CI2O -> NO2CI + OCI N0 2 + C102->N03 + 0C1 N0 2 + CO->NO + C02 Thus the description of nitrogen dioxide as a fairly reactive free radical is well illustrated by dimerization, combination with other free radicals560, addition, and abstraction reactions, with inorganic and organic molecules561. 13.7.2. Reactions of Liquid Dinitrogen Tetroxide533 Homolytic dissociation of N2O4 always occurs to a small extent in the liquid phase: N 2 0 4 (1) ^ 2N0 2 (1);
Δ/Zdiss = 19.5 kcal/mol
but of the various alternative heterolytic dissociation processes, it is perhaps surprising that the nitronium and nitrite ions have never been recognized as free ions in liquid N 2 0 4 . In fact the chemistry of the N2O4 solvent system is rationalized on the basis of self-ionization to give nitrosyl and nitrate ions according to the equation N 2 0 4 ^NO+ + N03 analogous to the ionization of water or liquid ammonia: 2H 2 O^H 3 0 + + OH-
558 j . H . T h o m a s , in Oxidation and Combustion Reviews, Vol. I (ed. C . F . H . Tipper), Elsevier, Amsterd a m (1965), p . 137. 559 A . F . T r o t m a n - D i c k e n s o n a n d G . S. Milne, Tables of Bimolecular Gas Reactions, N B S - 9 , W a s h i n g t o n (1969). 560 y . Rees a n d G . H . Williams, in Advances in Free-radical Chemistry, Vol. 3 (éd. G . H . Williams), L o g o s Press, L o n d o n (1969), p . 199. 561 A . V. Topchiev (translated by C. Matthews), Nitration of Hydrocarbons and other Organic Compounds, P e r g a m o n , L o n d o n (1959).
352
NITROGEN : K. JONES
Thus although the free ions in liquid N2O4 cannot be detected by methods such as infrared or Raman spectroscopy due to their low concentration, the large difference between the molar polarization and molar refraction is attributed562 to the presence of polar constituents in the liquid phase which are not present in either solid or gaseous states. These polar constituents are likely to be ion pairs rather than free ions because of the low dielectric constant of the solvent, and although ionic reactions will still occur, they do so as a result of collisions between ion pairs rather than free ions. Consequently, simple metal salts are insoluble in pure liquid N 2 0 4 , but do dissolve in mixtures of N2O4 diluted with solvents of higher dielectric constant533. Thus the addition of nitromethane (ε = 37) produces conducting solutions in which the self-ionization of N2O4 is enhanced to such an extent that dissociation is complete in solvents like perchloric or sulphuric acids: N2O4 + 3H2S04 -> NO+HS04 + HNO3 + SO3 + H30+HS04 Likewise a similar situation can be achieved by the addition of a donor solvent (such as ethyl acetate, diethyl ether, diethylnitrosamine, etc.). This arises from the electron-deficient nature of N2O4 and its tendency to form molecular addition compounds with donor molecules533, in particular oxygen and nitrogen bases bonded to the nitrogen(s) of N2O4: nD + N204 ^ D„ · N2O4 ^ (D„NO)+NO 3The wide variety of reactions in which liquid N2O4 can participate are summarized briefly below, although it should not be assumed that all such reactions are necessarily unique to the liquid phase. Acid-base reactions are described by the neutralization equation NO+ + N O i ^ N 2 0 4 Thus in liquid N2O4, nitrosyl chloride reacts with silver nitrate 563 : NOC1 + AgN0 3
N2o4
> AgCl + N 2 0 4
comparable with the corresponding reactions in water or liquid ammonia: HC1 + KOH H2°> KC1 + H20 NH4CI + NaNH2 —-i> NaCl + 2NH3 Similarly, compounds that produce "basic" solutions in liquid N2O4 are typified by alkylammonium nitrates which introduce excess nitrate into the system. Thus the reaction of zinc in these solutions illustrates the amphoteric behaviour of zinc compounds564 in N204: Ζη + 2ΕίΝΗ 3 ·Νθ3 + 2Ν 2 θ4 -* [EtNH 3 ] 2 [Zn(N0 3 )4] + 2NO
again directly analogous to reactions in water and liquid ammonia: Zn + 2NaOH + 2H 2 0 -> Na2[Zn(OH)4] + H 2 Zn + 2NaNH 2 +2NH 3 -> Na 2 [Zn(NH 2 )4]+H2 5
62 C. C. Addison, H. C. Bolton and J. Lewis, / . Chem. Soc. (1951) 1294. « C. C. Addison and R. Thompson, / . Chem. Soc. (1949) S211. 564 c . C. Addison and N. Hodge, / . Chem. Soc. (1954) 1138. 5
CHEMISTRY AND CHEMICAL PROPERTIES OF N 2 0 4
353
In contrast, the nitrosonium compounds are the "solvo-acids" and increase the concentrat i o n ^ of the NO + ion in liquid N2O4. Thus using tin as example, Ν2θ4
Sn+2NOC1
► SnCl2 + 2NO
again analogous to the corresponding reactions in water and liquid ammonia: Sn + 2HC1 - ^ - > SnCl2 + H 2 Sn + 2NH4C1
NH
% SnCl2 + 2NH 3 + H 2
Silver, zinc, mercury and the alkali metals react quite readily when brought into contact with liquid N2O4. For example, sodium requires continual stirring to expose clean metal surface if the reaction is to proceed to completion : Na + N204 -> NaN03 + NO At room temperature nitric oxide is evolved, but at lower temperatures or under pressure the formation of blue-green N2O3 is apparent533. The reaction is again directly analogous to those in water and liquid ammonia: 2Na+2H20 -> 2NaOH + H2 2Na + 2NH3 -» 2NaNH2 + H2 However, a number of other metals (Cd, Mn, Co, Cu, In, U) are only soluble if the liquid N2O4 is diluted with a solvent of higher dielectric constant such as nitromethane. The mixture becomes more reactive by enhancement of the self-ionization process of N2O4: MeN02
Cu + 3N204 > Cu(N03)2 · N204 + 2NO Similarly, the presence of donor solvents like ethyl acetate, diethyl ether, dimethyl sulphoxide, etc., in liquid N2O4 form the basis of excellent solvent systems for metals giving a variety of products including some stable solvate compounds containing dinitrogen tetroxide of crystallizations33: Ca+2N204 Mg+3N204 Bi+3N204 Cu + N 2 0 4
EtOAc EtOAc Me 2 SO AcOH 5-30% PhCN
> Ca(N0 3 ) 2 + 2NO > Mg(N03)2-N204+2NO > Bi(N0 3 ) 3 -3Me 2 SO + 3NO > Cu(OAc)2
Cu+5N204 > Cu(N03) · 2PhCN · 4N204 + NO Solvolysis reactions in liquid N2O4 are reactions of a salt MX with the solvent in which the anion X is totally or partially replaced by the nitrate ion, being the anion characteristic of the medium : MX + N2O4 -> MN03 + NOX Many such reactions have been reported including the following selection533: R4NC14- N 2 0 4 ^ R 4 N · N 0 3 + NOC1 ZnCl2 + N 2 0 4 ^ Zn(N0 3 ) 2 + 2NOC1 CaO + 2 N 2 0 4 -* Ca(N0 3 ) 2 + N 2 0 3 N a 2 0 2 + N 2 0 4 -► 2NaN0 3 N a O H + N 2 0 4 -> N a N 0 3 + H N 0 2
354
NITROGEN : K. JONES
This reaction provides an excellent route to anhydrous metal nitrates from the corresponding halides, particularly where X = Br or I as NOX decomposes preventing the possible formation of nitrosyl compounds533: Til4 + 4N204 ^ Ti(N03)4+4NO + 2I2 The oxidation properties of N2O4 towards inorganic compounds generally involve the donation of oxygen atoms rather than the removal of electrons from metal ions. This is well illustrated by the unusual salts that have been prepared by the oxidation of sodium hyponitrite565 with N 2 0 4 : Na2N202
rapid
> £-Na 2 N 2 0 3
slow
> Na 2 N 2 O s
slow
» Na2N206
100°
ΝΗ 2 ΟΗΊ EtN0 3 > -> a-Na 2 N 2 0 3
rapid
> Na2N204
slow
> Na 2 N 2 0 6
NaOEtJ Accordingly, the reaction of N2O4 with cobalt forms cobalt(II) rather than cobalt(III) compounds, and its selective oxidizing properties in organic chemistry566 are unique, particularly its ability to restrict oxidation to a single stage when further oxidation would appear to be possible. Dinitrogen tetroxide reacts readily with water according to the overall equation 3 N 2 0 4 + 2H 2 0 -> 4 H N 0 3 + 2NO
Several mechanisms, more than one of which may be applicable, have been proposed to account for the absorption stage in nitric acid production536 (see section 18.3) including the following: N 2 0 4 (g) -> N2O4O) N 2 0 4 (1) + H20(1) -> HN0 3 (1) + HN0 2 (1) 2HN02(1) -> HN0 3 (l) + NO(l) N 2 0 4 - > N O + + NOä N 2 0 4 + H 2 0 -> H + + N 0 J + H N 0 2 2 H N 0 2 -> H 2 0 + NO + N 0 2
However, in spite of this ready reaction, simple hydrated salts do not dehydrate completely when immersed in liquid N 2 0 4 (e.g. MgS0 4 -7H 2 0 -+ MgS0 4 -2H 2 0), and the selective reactivity may be related to possible differences in the bonding of the water567. Metal carbonyls react with N2O4 under mild conditions and provide simple preparative routes to metal nitrates and their N 2 0 4 solvates533: Co2(CO)8 + 8 N 2 0 4 -> 2Co(N0 3 ) 2 · 2 N 2 0 4 + 8CO + 4NO Fe(CO)5 + 4 N 2 0 4 -> Fe(N0 3 ) 3 · N 2 0 4 + 5CO + 3NO Mn2(CO)io + N 2 0 4 -> Mn(CO) 5 N0 3 + [Mn(CO)x(NO)y] 5
« C. C. Addison, G. A. Gamlen and R. Thompson, / . Chem. Soc. (1952) 338.
566 B. O. Field and J. Grundy, J. Chem. Soc. (1955) 1110. 567 C. C. Addison and D . J. Chapman, / . Chem. Soc. (1965) 819.
355
ANALYSIS OF N 0 2 - N 2 0 4
The reaction of N2O4 with organometallic compounds like tetramethyltin is explosively violent even at —80° and the reagents require dilution with an inert solvent5*^.· Me4Sn + 2 N 2 0 4
EtOAc > Me2Sn(N03)2
Finally, the reaction of N2O4 with the electron acceptor molecule boron trifluoride has received much attention. The reagents react both in 1:1 and 1:2 ratio to give stoichiometric solid products which have been formulated as containing the nitronium ion 567 : N2O4
BF3
> NOJ[BF3N02]-
BF3
> NOÎ[F3BONOBF3]-
Thus the overall scheme for N2O4 dissociation can be considered as in the following scheme533 : 2 N 0 2 ^ N 2 0 4 ^ [ N 0 2 + + NO7] ^ ™
» NO+ + NO3donor solvents
BF,
N O j [BF3-N02]"
13.8.
BIOLOGICAL
DnN204
(D n NO) + N0 3 -
PROPERTIES570
The toxic properties of nitrogen dioxide are well known, but the presence of dangerous concentrations in everyday situations is not widely appreciated. NO2 is the most dangerous component of "nitrous fumes", the name given to mixtures of oxides of nitrogen of variable composition. These mixtures can arise as a result of many industrial processes involving combustion but particularly electric-arc processes, electric-arc welding and oxy-acetylene welding. Hazards occur when such operations are carried out in confined, unventilated spaces. It is not considered safe to work for prolonged periods in an atmosphere containing >5ppm NO2, and short periods at higher concentrations can lead to symptoms such as inflammation of the lower respiratory tract and a marked elevation in the blood platelet count. Continued inhalation of concentrations > 100 ppm has led to death. 13.9.
ANALYSIS OF
N02-N204571
In large enough amounts, nitrogen dioxide is readily recognized by its characteristic red-brown colour. It can also be identified in equilibrium with its dimer by spectroscopic methods552; NO2 gives infrared absorption at 1626 cm -1 . Ultraviolet spectroscopy can also be used because dinitrogen tetroxide absorbs up to about 390 πιμ, and nitrogen dioxide in the 390-500 ιημ region. Nitrogen dioxide gives a positive test with Griess reagent due 568 c . C . A d d i s o n , W . B . Simpson a n d A . Walker, / . Chem. Soc. (1964) 2360. 569 R . w . Sprague, A . B . G a r r e t t a n d H . H . Sisler, / . Am. Chem. Soc. 82 (1960) 1059. 570 F. Call, in Mellows Comprehensive Treatise on Inorganic and Theoretical Chemistry, Vol. 8, Suppl. I, Nitrogen, P a r t I, L o n g m a n s , L o n d o n (1964), p p . 604-9. 571 K. I. Vasu, in Mellows Comprehensive Treatise on Inorganic and Theoretical Chemistry, Vol. 8, Suppl. I I , Nitrogen, P a r t I I , L o n g m a n s , L o n d o n (1967), p . 632. 572 M . J. T a r a s , in Colorimetric Determination of Nonmetals (ed. D . F . Boltz), Interscience, N e w Y o r k (1958).
356
NITROGEN: K. JONES
to the presence of nitrite in solution, which can also form the basis of a colorimetric method of analysis572. The application of titrimetric methods is somewhat similar to those described for nitric oxide. Nitrogen dioxide can be oxidized with excess cerium (IV) which is back-titrated with sodium oxalate: N 0 2 + Ce4+ + H 2 0 -> N O " + Ce3+ + 2 H +
Reference should also be made, however, to the analysis of nitrous fumes. This usually refers to mixtures of nitric oxide, nitrogen dioxide and other oxides that exist in equilibrium with them (principally N2O3 and N2O4). Their toxicity has directed considerable attention to the detection and estimation of small amounts of such materials573. Thus in this connection, and frequently in other situations, only the total nitrogen oxide content is required, and this can be obtained by oxidation with alkaline hydrogen peroxide to nitrate which can be determined colorimetrically by the phenoldisulphonic acid method. This can be preceded by absorption in Griess-Ilosvay reagent to obtain a separate estimation of N 0 2 content if required. For full analysis of such nitrogen oxide mixtures, mass spectrometry provides the quickest most reliable method. However, more recently, gas Chromatographie techniques have been developed which are capable of separating mixtures containing oxides of both nitrogen and carbon as well as elemental nitrogen using helium as carrier gas and thermal conductivity cell detectors574.
14. DINITROGEN PENTOXIDE 14.1. PREPARATION
Dinitrogen pentoxide is usually prepared in the laboratory by the dehydration of concentrated nitric acid with phosphorus pentoxide 575 · 576 : 2 H N O 3 + P 2 O 5 -> N2O5 + 2HPO3
Considerable care must be exercised during this reaction, as the procedure577 involves either the dropwise addition of concentrated HNO3 (d 1.525g/cm3) onto a large excess of P2O5 at — 10°C followed by slow distillation at about 35°C, or by adding P2O5 in one portion to nitric acid previously cooled to — 78 °C then allowing the mixture to attain room temperature slowly578. The principal alternative method is by the oxidation of nitrogen dioxide with ozone: 2 Ν 0 2 + θ 3 - * Ν 2 θ 5 + θ2 573 Methods for the Detection of Toxic Substances in Air, Booklet N o . 5, Nitrous Fumes, Ministry o f Employment and Productivity, H M Factory Inspectorate, H M S O , London (1969). 574 j . Janak and M. Rusek, Chemické Listy 48 (1954) 397. 5751. R. Beattie, in ref. 571, pp. 269-77. 576 p . w . Schenk, in Handbook of Preparative Inorganic Chemistry, Vol. I (ed. G. Brauer), Academic Press, N e w York (1963), p. 489. 577 N . S. Gruenhut, M. Goldfrank, M. L. Cushing and G. V. Caesar, Inorganic Syntheses, Vol. I l l (ed. L. F. Andrieth), McGraw-Hill, N e w York (1950), p. 78. ^78 G. V. Caesar and M. Goldfrank, / . Am. Chem. Soc. 68 (1962) 372.