NITRIC OXIDE

323

ANALYSIS OF NITROUS OXIDE 10.9.

A N A L Y S I S OF N I T R O U S

OXIDE459

Thermal decomposition of nitrous oxide in the presence of platinum or palladium catalyst will liberate oxygen which can be detected with alkaline pyrogallol reagent, but is not specific to this particular oxide. The mass spectrometer is probably the most reliable technique for detecting small quantities of gas. However, infrared absorption spectroscopy is a more accessible technique, and depending on other gases present, the intensity of one of the infrared bands at 590, 1285 and 2224 cm - 1 can be used to estimate the nitrous oxide content of gas mixture. In connection with the use of nitrous oxide as an anaesthetic, its estimation in admixture with oxygen and ether can be carried out by the acoustic gas analyser which, as its name implies, is based on the variation of the velocity of sound according to the density of the gas. Hence by pre-calibrating for various mixtures, nitrous oxide content can be continuously monitored. If nitrogen is also present, oxygen has to be determined separately by means of an oxygen analyser for the four components to be estimated. Combustion methods include reduction over a palladium catalyst followed by volumetric determination of the nitrogen : N20 + H 2 ^ N 2 + H20

Alternatively, nitrous oxide can be reduced over heated copper gauze: N 2 0 + C u - > C u O + N2

or as another variation to eliminate interference from carbon dioxide: N20 + CO-^N 2 + C02

11. N I T R I C OXIDE 11.1. P R O D U C T I O N AND INDUSTRIAL

USE

From an industrial point of view, nitric oxide is probably the most important oxide of nitrogen460. It is produced by the catalytic oxidation of ammonia461»462 at 800-960°C: 4NH 3 + 5 0 2 - ^ : ^ 4 N O + 6 H 2 0 ;

Δ#(298) = -216.55 kcal/mol

This, and the alternative reactions which are possible in ammonia oxidation, have been considered in more detail in section 3.10.2 The direct reaction of nitrogen and oxygen requires high temperatures before equilibrium conditions are such as to bring any appreciable yields of nitric oxide. The BirkelandEyde electric arc oxidation of atmospheric nitrogen was developed during the 1920's in countries where hydroelectric power was cheap, but the process is now obsolete460. More recently the combustion of air-natural gas mixtures in the "pebble-bed" or Wisconsin process can produce temperatures of over 2000°C, and although the process is feasible, it is not at present economic463. The equilibrium concentrations of NO for each process 459 K. I. Vasu, in Mellows Comprehensive Treatise on Inorganic and Theoretical Chemistry, II, Nitrogen, Part 2, Longmans, London (1967), p. 628. 460 i . R . Beattie, in ref. 459, Section X X I I , p . 158. 461 I. R . Beattie, in ref. 459, Section X X V , p . 216. 462 G . C . B o n d , Catalysis by Metals, A c a d e m i c Press, N e w Y o r k (1962), p . 456. 463 E . D . E r m e n c , Chem. Engng. Progress 52 (1956) 149.

Vol. 8, Suppl.

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NITROGEN: K. JONES

are 4 % and 2 % respectively, and the gas mixture has even to be cooled rapidly to prevent any further loss by decomposition (see section 11.7.1). Clearly, the fate of most nitric oxide is oxidation to nitrogen dioxide almost as soon as it is formed, and the principal industrial application of both oxides is as intermediate oxidation products in the manufacture of nitric acid464. The presence of nitric oxide in the atmosphere can obviously occur through lightning discharge and it is also present in automobile exhausts465, hence subsequent oxidation and dissolution produces nitric acid in very low concentration in rain water. However, although this could be considered as providing a beneficial source of fixed nitrogen, the build-up of concentrations of nitrous fumes in city atmospheres is likely to be a considerable health hazard. Nitric oxide is also present in the upper atmosphere, and important in photochemical reactions arising from solar radiation466. 11.2. LABORATORY P R E P A R A T I O N A N D P U R I F I C A T I O N OF N I T R I C O X I D E S ?

Nitric oxide may be prepared from aqueous solution by the reaction between nitrous acid and iodine or ferrocyanide : 2 K N 0 2 + 2KI + 2H 2 S0 4 -> 2NO + 2 K 2 S 0 4 + 1 2 + 2H 2 0 K N 0 2 + K 4 Fe(CN) 6 + 2CH3COOH -> NO + K 3 Fe(CN) 6 + 2CH3COOK + H 2 0

When dilute sulphuric acid is added dropwise on to sodium nitrite, the nitrous acid formed initially decomposes to give a steady stream of nitric oxide: 6NaN02 + 3H2S04 -> 4NO + 2H20 + 3Na2S04+2HN03 The gas is purified by passing through 90 % sulphuric acid followed by 50 % potassium hydroxide then dried by passing through a trap cooled in solid C0 2 -ether mixture, and condensed over phosphorus pentoxide cooled with liquid nitrogen. On fractional distillation, the middle cut is collected and redistilled to give pure NO. Dry nitric oxide can be obtained directly simply by heating the solid mixture of chromium^ 11) oxide with nitrate and nitrite, preferably in the presence of calcined iron(III) oxide468 : 3KN02 + KN03 + Cr203 -> 4NO + 2K2Cr04 Nitric oxide is also formed in conjunction with other gases during the reduction461 of nitric acid, nitrates and nitrites, or nitrosyl derivatives: 8HN0 3 + 3Cu

> 2NO + 3Cu(N0 3 ) 2 + 4H 2 0

2Ba(N0 2 ) 2 +1 2 - ^ > 2NO + Ba(N0 3 ) 2 + BaI2 2NOC1 + 21-

> 2NO + 1 2 + 2X~

464 C. J. Pratt and R. Noyes, Nitrogen Fertilizer Chemical Process, Noyes Development Corporation, New Jersey (1965). 465 W. B. Innes and K. Tsu, in Kirk-Othmer Encyclopedia of Chemical Technology, 2nd edn., Vol. 2, Interscience, New York (1963), p. 814. 466 J. Heicklen and N. Cohen, in Advances in Photochemistry, Vol. 5 (eds. W. A. Noyes, Jr., G. S. Hammond and J. N. Pitts), Interscience, New York (1968), pp. 157-328. 467 P. W. Schenk, in Handbook of Preparative Inorganic Chemistry, Vol. I (ed. G. Brauer), Academic Press, New York (1963), p. 460. «* R. A. Ogg, Jr. and J. D. Ray, / . Am. Chem. Soc. 78 (1956) 7993.

NUCLEAR ISOTOPES OF NITRIC OXIDE

325

Nitrosonium salts react with hydroxylic solvents: NO++OH- ^ HN0 2 F^ NOi + H+ Thus the equilibrium is only displaced to the left in strongly acidic media. Nitrosonium hydrogen sulphate is formed therefore by dissolution of nitrites or dinitrogen trioxide in sulphuric acid 469 : N2O3 + H2SO4 -> 2NO++3HS04 + H 3 0 + although hydrolysis occurs in solutions containing less than 80% H2SO4.

11.3. NUCLEAR ISOTOPES OF NITRIC OXIDE470 15

NO is commercially available or can be prepared by the reduction of nitrogen dioxide, nitrites or nitrates with mercury and sulphuric acid: 6i5N0 2 +4Hg+2H 2 S0 4 -> 4i5NO + Hg2(N03)2 + 2HgS04+2H20 Conversion of ammonia into nitric oxide can be achieved most conveniently without proceeding via nitrogen by direct oxidation with permanganate in an autoclave for 7-8 hr at 170-180°C to nitrate which can be reduced to nitric oxide as before. By passing a mixture of oxygen and nitrogen through a high-tension electric arc, a mixture of N2O3 and N2O4 is formed which can also be reduced to NO as above. This method has been used in the preparation of 1 5 NO and N ^ O by using the appropriate enriched elements 471 .

11.4. STRUCTURE AND BONDING 472

X-ray crystallography shows that solid nitric oxide consists of a loosely bound dimeric (NO)2 species which may have a rectangular shape. The residual entropy and other physical properties461 also suggest that complete randomization of the two orientations occurs, and although the mean distance between the two NO units is about 2.4 Â, it is not possible to tell whether they are oriented in the same direction or head to tail as illustrated in the structure: N O I I 1.10 A O N 2.38 À The blue-coloured liquid phase is also asssociated and infrared studies 473 provide some evidence for N - N bonds implying a cw-configuration with the structure: •N—frill II O O However, nitric oxide is normally encountered as a gas, in which phase it is colourless and monomeric with little tendency to associate via electron pairing. 469 c . C. Addison and J. Lewis, Quart. Rev. (9) (1955) 115. 470 w . Spindel, in Inorganic Isotopic Synthesis (ed. R. H. Herber), Benjamin, New York (1962), p. 100. 471 I. Dostrovsky and D. Sammuel, in Inorganic Isotopic Synthesis, (ed. R. H. Herber), Benjamin, New York (1962). 472 w . J. Dulmange, E. A. Meyers and W. N. Lipscomb, Acta Cryst. 6 (1953) 760. 473 w . G. Fateley, H. A. Bent and B. Crawford, / . Chem. Phys. 31 (1959) 204.

326

NITROGEN : K. JONES

Using molecular orbital theory, the electronic structure of nitric oxide in the ground state is usually written474 NO 2Π (1σ+)2(2σ+)2(3σ+)2(1π)4(4σ+)2(5σ+)2(2π)ΐ

Molecular orbital calculations suggest475 that the 3σ, 4σ and 5σ orbitals are bonding, nonbonding and anti-bonding respectively, with the last electron entering into an antibonding π orbital and accounting for the paramagnetism. Furthermore this theory predicts the bond order to be 2.5, and that it should not be too difficult to remove an electron to form the NO+ ion with a shorter and stronger bond than NO itself. In fact, each of these predictions is correct, the bond length of NO (1.14 Â) lies between that of a double (1.18 Â) and a triple bond (1.06 Â), the ionization potential of 9.25 eV is appreciably lower than for similar molecules (N2, 0 2 and CO are 15.6, 12.1 and 14.0 eV respectively) and the stretching frequency of the NO + ion in nitrosyl salts (2150-2400 cm -1 ) is higher than nitric oxide itself (1888 cm"1). In fact, salts of the nitrosyl cation, or the nitrosonium ion as it is alternatively named, such as the perchlorate, bisulphate and tetrafluoroborate, are well-characterized crystalline solids which are isomorphous with corresponding hydroxonium and ammonium salts 4 ^. The NO+ species is also isoelectronic with N 2 , CO and CN _ , each of which can be considered by molecular orbital theory as having an electronic structure containing no antibonding π electrons and a bond order of 3. Thus it is not surprising that a wide range and variety of NO complexes of transition metals (see Chapter 46) analogous to metal carbonyls are known. In most cases, NO transfers an electron to the metal, and the NO+ so formed further donates a lone pair of electrons from nitrogen to the metal to give compounds like Fe(NO)2(CO)2 and Co(NO) (CO)3 which are stable and isoelectronic with Ni(CO)4. The valence bond approach is also worthy of mention, for although NO is often represented by the resonance forms + - + * N—O <-> N = 0 <-► N = 0 · <-+ N = 0 *

it is adequately described by the Linnett non-pairing structure474: :N=0: +

while the cation NO is [:N^O:]+

Similarly, the anion NO~ would on the same basis have the structure [:N=0:]-

however, most compounds in which these ions were thought to be present are now known to contain the anion N 2 0|~ 476. 474 M. Green, Developments in Inorganic Nitrogen Chemistry, Vol. I (ed. C. B. Colburn), Elsevier, Amsterdam (1966), p. 28. 47 s H. Brion, C. Moser and M. Yamazaki, / . Chem. Phys. 30 (1959) 673. 47Ö N. Gee, D. Nicholls and V. Vincent, / . Chem. Soc. (1964) 5897.

MOLECULAR CONSTANTS OF NITRIC OXIDE

327

11.5. MOLECULAR CONSTANTS Various molecular constants of nitric oxide are listed in Table 67. The electronic absorption spectrum of NO extends from just below 2300 Â to the vacuum ultraviolet region. Transitions between the ground electronic state and the Α2Σ+, B2Ut9 C2U and Ό2Σ+ states are usually referred to as y, β, δ and ε bands respectively and have been extensively investigated466'4?7. Infrared and Raman spectra of various isotope isomers of nitric oxide as gas, liquid, solid and matrix trapped molecule samples also indicate the presence of eis or trans dimer (ONNO) forms47?. However, the fundamental bands are also split due to an interaction between the angular and spin momentum vectors (L = 1, S = i) which may couple L + S o r L- S dividing the 2 Π ground state into

TABLE 67. MOLECULAR CONSTANTS OF NITRIC OXIDE3

Property X-ray diffraction data5 System Space group Lattice constants N-O bond length N · · · O intermolecular distance Electron diffraction datac N-O bond length Fundamental vibration frequency Force constant Moment of inertia Dipole moment Bond dissociation energy First ionization potential Magnetic susceptibility (293 K) Magnetic moment (296 K)

Value

Monoclinic P2i/a a = 6.68, b = 3.96, c = 6.55, a = 127.9° 1.10Â 2.38 Â 1.15 A 1876 cm-i 1.59 x 106 dynes/cm . 16.47 x 10-40 g / c m 2 0.15 D 6.50 eV, 149.9 kcal 9.25 eV 1.46 x 10 _3 c.g.s. units 1.837 BM

a

Ref. 461. Ref. 472. L. E. Sutton, Tables of Interatomic Distances and Configurations in Molecules and Ions, Chemical Society, London (1958). b c

an upper paramagnetic state 2Π3/2 and a lower diamagnetic state 2 Π|. Interaction with the rotational motion of the molecule also causes the upper state to split into four sub-levels giving rise to three absorption lines with the usual selection rule ms = ± 1 , and in fact they are each observed to be further split into three components due to the interaction of the spin of the nitrogen nucleus461. The energy difference between the paramagnetic and diamagnetic 2Π states is only about 350 cal/mol and as kt is considerably greater than this value at room temperature and comparable down to even quite low temperatures, it is this 477 R . F . Barrow and A. J. Merer, in ref. 459, p. 507.

328

NITROGEN: K. JONES

phenomenon that gives rise to the temperature dependence of the magnetic moment. The hyperfine splitting of the electron spin resonance spectrum of NO is also explained on the basis of a strong interaction between the moment of the unpaired electron and the rotation of the molecule, and also leads to the conclusion that about 60% of the spin density is concentrated on the nitrogen atom478. Microwave spectral data have been used to obtain values for magnetic hyperfine structure constants, the nuclear quadrupole resonance moment and the dipole moment477. The various dissociative ionization processes have also been studied in great detail and their corresponding appearance potentials are consistent with the value of 6.50 + 0.01 eV for the dissociation energy of nitric oxide. Photo-ionization and electron impact studies are in agreement that the ionization potential for the formation of NO + is 9.25 + 0.02 eV20. 11.6. PHYSICAL PROPERTIES

Physical constants for nitric oxide are listed in Table 68. 11.7. CHEMISTRY AND CHEMICAL PROPERTIES 11.7.1. Decomposition of Nitric Oxide As the free energy change in the formation of nitric oxide from the elements is large and positive: N 2 + 0 2 -> 2NO ;

Δ G = + 20 kcal/mol

the isolation of a reasonable yield of product requires not only a high temperature but also a rapid rate of cooling of the equilibrium mixture. The equilibrium constant of the reaction K = [NO]2/[N2][02], which at 298 K is 5.27 x 10 -31 , has been calculated for a wide range of conditions, but their experimental confirmation has been difficult. Although many products have been reported for the thermal decomposition of nitric oxide, the principal reaction to be considered is that reverting to the elements: 2NO-+N2+O2

The results of several investigations indicate that at 1130-1330°C the decomposition is described by a second-order homogeneous reaction, whilst at higher temperatures reactions involving atomic nitrogen and nitrogen dioxide become important461. Heterogeneous decomposition of nitric oxide on platinum and platinum-rhodium wire above 1000°C is bimoleculâr with respect to NO and is retarded proportionally to the oxygen concentration. Photochemical decomposition of nitric oxide has also been widely investigated, and the products vary to some extent with the state of excitation from which the nitric oxide is decomposed466. 11.7.2. Oxidation Undoubtedly the most important reaction of nitric oxide is its oxidation to nitrogen dioxide460: 2NO + 0 2 - > 2 N 0 2 478 R . Beringer, E. Rawson and A. Henry, Phys. Rev. 94 (1954) 343. 479 G . R. A. Johnson, in ref. 459, p. 544.

329

PHYSICAL PROPERTIES OF NITRIC OXIDE TABLE 68. PHYSICAL PROPERTIES OF NITRIC OXIDE

Temperature (K)

Property Melting point Boiling point Critical temperature Critical pressure Critical density Affusion AH vaporization /\Hf (enthalpy of formation)" AHf (enthalpy of formation) Δ Gf (Gibbs energy of formation) S° (entropy) C°p (heat capacity) Density of solid (calculated from X-ray) Density of liquid

Density of gas Viscosity of liquid Thermal conductivity Virial coefficient0 B (PV/nRT) = l+nB(T)+ V

...

109.49 121.36 179.15 109.49 121.36 0 298 298 298 298 20 78 98 110.15 113.65 117.15 119.55 293 120 200 280 120 300 121.72 153.06 274.00 277.60 310.94

Value -163.6°C -151.8°C -94°C 65atm 0.52 g/cm3 0.5495 kcal/mol 3.292 kcal/mol 21.45 kcal/mol 21.57 kcal/mol 20.69 kcal/mol 50.347 cal/deg mol 7.133 cal/deg mol 1.57g/cm3 1.556 g/cm3 1.46 g/cm3 1.332 g/cm3 1.306 g/cm3 1.277 g/cm3 1.227 g/cm3 1.3402 g/1 843.6x107 g/cm s 1371.2x107 g/cm s 1837.6x107 g/cm s 2.580 g cal/cm s 6.189 g cal/cm s -224.4 -107.9 -22.8 - 2 6 . 2 ( c = 1400) - 1 9 . 0 ( c = 1900)

Vapour pressure of solid P log Ρςι„ = - ^ | ^ + 0.00076Γ+ 9.05125 Vapour pressure of liquid P logPcm =

Ζ

^+0.002364Γ+8.562128

K

°C

P

K

°C

P

88.6 94.9 101.4 107.1 116.3

-184.55 -178.25 -171.75 -166.05 -156.85

1 mm 10 mm 40 mm 100 mm 400 mm

121.4 128.0 137.4 146.8

-151.75 -145.1 -135.7 -127.3

1 atm 2atm 5 atm 10 atm

a

Ref. 461. D. D. Wagman, W. H. Evans, V. B. Parker, I. Halow, S. M. Bailey and R, H. Schumm, Selected Values of Chemical Thermodynamic Properties, NBS Technical Note 270-3, Washington (1968). c J. H. Dymond and E. B. Smith, The Virial Coefficient of Gases—A Critical Compilation, Oxford (1969). b

330

NITROGEN: K. JONES

Work on this classical termolecular reaction was first carried out by Bodenstein480 in 1918, and extensive investigations which have been reported since that date confirm that it follows a simple third-order rate law over a wide range of experimental conditions : -d(NO)ldt

= *[NOP[0 2 J

However, a further unusual feature of this reaction is that the rate constant decreases with increasing temperature. This can be accounted for in the mechanism which postulates the initial formation of dimer : 2NO ^ N 2 0 2 Ν 2 θ2 + θ 2 ^ 2 Ν 0 2

Since the equilibrium constant K for the dimerization step should decrease with increasing temperature, the process can now be described by the rate law: -έ/[ΝΟ]/Λ = *'#[ΝΟΡ[θ2]

where K is the equilibrium constant for the dimerization and k'K is synonymous with k. Thus the negative temperature coefficient is explained on the basis that since K should be expected to decrease with temperature, it is not unreasonable to expect k'K to do the same, particularly with a very low energy of activation. An alternative mechanism481 involves the species NO3, together with the possibility of the pernitrite-nitric oxide complex. Like nitrogen trioxide, dinitrogen pentoxide has also been detected by infrared spectroscopy482 during nitric oxide-oxygen reactions which may lend support for the scheme : NO+02^N03 N0 3 + NO->2N02 N0 3 + N0 2 ->N 2 0 5 N205 + NO->3N02 11.7.3. Reactions of Nitric Oxide Being the simplest stable molecule with an odd number of electrons, nitric oxide has been well investigated with regard to its reactivity towards atoms, free radicals and other paramagnetic species. Many of these reactions have provided unique systems for study by gas kinetics466»483 and the results of which have led to an understanding of their mechanisms. The use of nitric oxide in the monitoring of atomic nitrogen has already been mentioned in section 2.12, while chemiluminescence due to the reaction NO + 0-^N0 2 +Av has likewise been used for both the qualitative and quantitative estimation of atomic oxygen484. The termolecular reaction NO+0+M->N0 2 +M is also well known and first order in each reactant. 480 M . Bodenstein, Helv. Chim. Acta 18 (1935) 743. 481 J. D . Ray and R. A . Ogg, Jr., / . Chem. Phys. 2 6 (1957) 984. 482 R. A . Ogg, Jr., / . Chem. Phys. 18 (1950) 770. 483 A . F. Trotman-Dickenson and G. S. Milne, Tables of Bimolecular Gas Reactions, National Standard Reference Data Series in National Bureau o f Standards, 9, Washington (1969). 484 F. Kaufman, in Progress in Reaction Kinetics, Vol. I (ed. G. Porter), Pergamon, Oxford (1961), p. 1.

CHEMISTRY AND CHEMICAL PROPERTIES OF NITRIC OXIDE

331

Of similar major importance is the reaction of NO with hydrogen atoms produced by electric discharge4»* to give nitroxyl HNO, which has also been prepared by the flash photolysis of nitroalkanes or isoamyl nitrite. Nitroxyl has also been postulated as an intermediate in several industrial processes (see section 3.10.2), and its infrared spectrum in an argon matrix at 20°C has been recorded4^, as well as the value of 48.6 kcal/mol for the bond strength^? Z)(H-NO). In the reactions of NO with the halogens, both atomic and molecular species need to be considered and their mechanisms are among those generalized as follows: NO + X 2 ^ X 2 N O NO + X 2 N O - ^ 2 X N O NO + X 2 - > X N O + X NO + X - > X N O * X N O * - ^ X N O + /zv XNO* + M - * X N O + M

Although the NO-F2 reaction gives FNO predominantly, some ONF3 can be obtained by photochemical fluorination of NO even at room temperature488. The nitric oxide-chlorine system has been most thoroughly investigated as both radical and molecular halogen reactions are important contributing factors in the overall mechanism466»489. Less is known about the NO-Br2 reaction, but with iodine the mechanism is thought466 to be NO + I^INO INO+I->NO + I2 Likewise, the reaction of NO with HI NO + 6HI -> N H 4 I + H 2 0 + 5/212

is thought to involve the following steps490: NO + H I - > H N O + I HNO + H I - > H 2 + I N O H N O + I 2 - > H I + INO INO + I ^ N O + I 2

Reactions with other halogen compounds are equally complex. Nitryl chloride reacts rapidly with NO 491 , but it is uncertain whether this reaction involves oxygen or chlorine transfer. NO+CINO2 -> N 0 2 4- C1NO

485 486 487 488 489 490 491

M . A . A . Clyne and B. A . Thrush, Trans. Faraday Soc. 57 (1961) 1305. H . W . Brown and G. C. Pimentel, J. Chem. Phys. 29 (1958) 883. M . J. Y . Clement and D . A . Ramsay, Can. J. Phys. 39 (1961) 205. E . W . Lawless and I. C. Smith, Inorganic High-energy Oxidizers, M . Dekker, N e w York (1968), p. 68. p . G. Ashmore and M. S. Spencer, Trans. Faraday Soc. 55 (1959) 1868. J. L. Holmes and E. V. Sundaram, Trans. Faraday Soc. 6 2 (1966) 910. E . C. Freiling, H. S. Johnston and R. A . Ogg, / . Chem. Phys. 20 (1952) 327.

332

NITROGEN: K. JONES

Nitric oxide and hydrogen chloride condense to a deep blue liquid which may contain the complex NOHCl. With xenon fluorides, the reactions with NO have been followed by mass spectrometry which indicates49* stepwise removal of fluorine: XeF 4 +NO -* XeF 3 + FNO XeF 2 + NO -> XeF+FNO

Nitric oxide can, however, be oxidized with iodine pentoxide461 : IONO+2I2O5 -► 2I2+5N2O4 IONO+3I2O5 -> 3I2+5N205 Halogenated methyl radicals add to NO to form nitroso compounds466 of which, perhaps, the best known493 is the blue gas trifluoronitrosomethane CF3NO: CF 3 + N O - * C F 3 N O CFCl2 + N O ^ C F C l 2 N O CF2H + NO -> CF 2 HNO -> CF 2 =NOH

However, with diboron tetrahalides breakdown of the unstable adducts occurs 494 » 4 ^: 3B 2 Cl 4 +3NO - ^ ^ 3B 2 Cl 4 NO -> B 2 (NO) 3 BCl 3 + BCl3

I B 2 (NO) 3 + BCl3 > 4BF 3 + 3N 2 0 + B 2 0 3 > 8BF3 + 3 N 2 + 2 B 2 0 3

3B2F4 + 6NO 6B 2 F 4 + 6NO

The primary step in the reaction of nitric oxide with ozone is a bimolecular secondorder reaction466: N 0 + 0 3 ^ NO2+O2;

Δ#(298) = - 4 7 . 8 kcal/mol

with possible secondary reactions: 2N02 + 0 3 -^N 2 0 5 + 0 2 2NO + 0 2 ^ 2 N 0 2 Similarly, the gas phase oxidation of NO with HNO3 also gives nitrogen dioxide496: NO + 2HN03 -> 3N02+H20 The reactions of NO with the oxides of sulphur have been investigated in view of their relevance to the chamber process for the manufacture of sulphuric acid497, and losses occur when nitrous oxide is produced : 2NO + S0 2 ->N 2 0 + S03 NO+2S03-*(S03)2NO 2NO + 2S02+2H20 -> 2H2S03NO -> H2S03(NO)2 -^ N20+H2S04 + H2S03 492 H. S. Johnston and R. Woolfolk, / . Chem. Phys. 41 (1964) 269. 4 « j . Banus, / . Chem. Soc. (1953) 3755. 494 A. K. Holliday and A. G. Massey, / . Inorg. Nucl. Chem. 18 (1961) 108. «>5 A. K. Holliday and F. B. Taylor, / . Chem. Soc. (1962) 2767.

4 96 J. H. Smith, / . Am. Chem. Soc. 69 (1947) 1741. 497 T. j . p # Pearce, in Inorganic Sulphur Chemistry (ed. G. Nickless), Elsevier, Amsterdam (1968), p. 543.

CHEMISTRY AND CHEMICAL PROPERTIES OF NITRIC OXIDE

333

The intermediate dinitrososulphites can be prepared as follows467: 2NO + K 2 S 0 3 -+· K 2 S0 3 (NO) 2 6NO + N a 2 S 2 0 4 + 2 N a O H -> 2 N 2 0 + N a 2 S 0 3 · N 2 0 2 + N a 2 S 0 4 + H 2 0

Nitrogen compounds which react with NO include ammonia as in the side-reactions during NO production by ammonia oxidation464: 2NH 3 + 3NO -> 5/2N2 + 3 H 2 0

NO at 20 atm pressure is also taken up by solutions of nitrosonium salts forming the deep blue N2OJ cation: which is thought to be responsible for the colour during the chamber process for the production of sulphuric acid498 : NO+HS04 + NO -> N 2 O J H S O ;

Nitrogen trichloride reacts with nitric oxide 4 " and at — 150°C the overall reaction is NCI3 + 3NO -> 2NOC1 + N 2 0 + Cl

but at — 80°C the main reaction is represented by NCI3 + 2NO -> N 2 0 + NOC1 + Cl 2

Reaction of nitric oxide with carbon compounds, in particular organic-free radicals, have been reviewed466. The explosive reaction with carbon disulphide has received much attention461 because of the associated light emission. The reaction of sodium oxide with NO gives sodium hydronitrite initially which then decomposes500: 4 N a 2 0 + 4NO

> 4 N a 2 N 0 2 -> 2Na 2 0 + 2 N a N 0 2 + N a 2 N 2 0 2

With alkali metal hydroxides, both N 2 0 and N2 are formed501 : 4NO + 2MOH -* N 2 0 + 2 M N 0 2 + H 2 0 6 N O + 4 M O H -+ N 2 + 4 M N 0 2 + 2 H 2 0

Sodium reacts with NO directly or in liquid ammonia to give (NaNO)*502; which on the evidence of its X-ray diffraction pattern is thought to be different to ordinary sodium hyponitrite and possibly may be α 5 - ^ 2 Ν 2 0 2 . Infrared spectra of the product of reaction of lithium with NO in an argon matrix suggests LiON rather than LiNO 503 . A solution of hyponitrous acid is obtained when a stream of NO is passed through an ether solution of lithium aluminium hydride504. 498 F . Seel, B . Ficke, L . Riehl a n d E. Vo'lkl, Z. Naturforsch. 86 (1953) 607. 499 w . A . N o y e s , / . Am. Chem. Soc. 53 (1931) 2137. 500 E . Zintl a n d H . H . B a u m b a c h , Z. anorg. allgem. Chem. 198 (1931) 88. soi E . Barnes, / . Chem. Soc. (1931) 2605. 502 H . Gehlen, Ber. 72B (1939) 159. 503 w . L . S. A n d r e w s a n d G . C . Pimentel, / . Chem. Phys. 44 (1966) 2361. 504 p . K a r r e r a n d R . Schwyzer, Rec. Trav. Chim. 69 (1950) 474.

334

NITROGEN: K. JONES

The reaction of nitric oxide with transition metal compounds^ i s an important route to nitrosyl complexes which will be discussed in Chapter 46. Classes of compounds containing the nitrosyl group including metal nitrosyls [Ru(NO)4], nitrosyl carbonyls [Co(NO)(CO)3], nitrosyl halides [Fe(NO)2I], nitrosyl pseudohalides [Fe(NO)(CO)5]2and organometallic (principally cyclopentadienyl) nitrosyl derivatives506 [(7r-C 5 H 5 )2Mn 2 (NO)3]

have been extensively investigated. 11.8. ANALYSIS OF NITRIC OXIDE The classic qualitative test for NO is the brown ring test which is demonstrated by the interaction of NO with aqueous iron(II) solution to give [Fe(NO)]2+. A standard quantitative estimation is based on the assumption that nitric and nitrous acid solutions produce pure NO when shaken with concentrated sulphuric acid which can be measured volumetrically in a Lunge nitrometer. Combustion methods based on the reaction of NO with carbon monoxide lead to the volumetric measurement of nitrogen after absorbing other residual gases507. 2NO + 2CO-*2C02+N2 Of the other chemical methods, NO can be absorbed directly in an excess of acid permanganate solution and the residual permanganate reduced with excess iron(II) solution which is then back-titrated with standard permanganate. 10NO + 6KMnO 4 +9H 2 SO 4 -+ 3 K 2 S 0 4 + 6 M n S 0 4 + 10HNO 3 +4H 2 O

Most other methods depend on the conversion of NO to nitrites: 4NO + 0 2 + 4KOH -> 4KN02+2H20 2NO+HN03 + 3H2S04 -> 3NOHS04

3H2

°> 3HN02+3H2S04

and subsequent determination either by titration with permanganate: 5HN02+2KMn04 + 3H2S04 -> K2S04+2MnS04+5HN03 + 3H20 or colorimetrically508 by a diazo coupling reaction such as Griess reagent (sulphanilic acid and a-naphthylamine). Of the physical techniques, infrared spectroscopy is applicable and the bands at 1850 and 1925 cm -1 useful to estimate the concentration of NO in the gas phase. However, mass spectrometry is probably the best method of determining NO content, and a good example of its potential is its ability to monitor NO concentrations in the exhausts of internal combustion engines509. 505 B. F. G. Johnson and J. A. McCleverty, in Progress in Inorganic Chemistry, Vol. 7 (ed. F. A. Cotton), Interscience, New York (1966), p. 277.

506 w . P. Griffith, in Advances in Organometallic Academic Press, N e w York (1968), p. 211. 507 K. I. Vasu, in ref. 459, pp. 563-673. 508

(1958). 509

Chemistry,

Vol. 7 (eds. F. G. A . Stone and R. West),

M. J. Taras, in Colorimetric Determination of Non-metals (éd. D. F. Boltz), Interscience, New York R. D. Craig, in Modern Aspects of Mass Spectroscopy (ed. R. I. Reed), Plenum, New York (1968).