Chapter 16

NITROGEN

16-1 Introduction The nitrogen atom (ls22s22Px2py2pz) can complete its valence shell in the fol­ lowing ways:

3 1. Electron gain to form the nitride ion N -; this ion is found only in saltlike nitrides of the most electropositive metals.

2. Formation of electron-pair bonds: (a) single bonds, as in NH3 , or (b) multiple bonds, as in :N=N:, -N=N-, or N02 • 2 3. Formation of electron-pair bonds with electron gain, as in NH2or NH -. 4. Formation of electron-pair bonds with electron loss, as in the tetrahedral

ammonium and substituted ammonium ions, [NR4 ]+. The following structural types (recall Chapter 3) are common among those

compounds of nitrogen having covalent bonds: AB 4 (as in tetrahedral NR:); AB 3E (as in pyramidal NR3 ); AB 2E2 (as in bent NR2); AB3 (as in planar NOg); AB 2E (as in bent R2 C=N-OH); and ABE (N2 ). There are a few stable species in which, formally, the nitrogen valence shell is incomplete. Nitroxides,

R2N=Q, NO, and N02 are the best examples; these have unpaired electrons and are paramagnetic.

Three-Covalent Nitrogen

The molecules NR3 are pyramidal; the bonding is best considered as involving sp3 hybrid orbitals so that the lone pair occupies the fourth position. There are three points to note:

1. As a result of the nonbonding electron pair, all NR3 compounds behave as Lewis bases and they give donor-acceptor complexes with Lewis acids,

for example, F3B:N(CH3h, and they act as ligands toward transition metal ions as in, for example, [Co(NH3 )6]3+. 2. Pyramidal molecules (NRR'R") should be chiral. Optical isomers cannot be isolated, however, because such molecules very rapidly undergo a mo­ tion known as inversion in which the N atom oscillates through the plane of the three R groups, much as an umbrella can turn inside out (Fig. 16-1). The energy barrier for this process is only about 24 kJ mol-I. 3. There are a very few cases where three-covalent nitrogen is planar; in

these cases multiple bonding is involved as we discussed for N (SiMe3 ) 3' in 399 400 Chapter 16 / Nitrogen

Q' N R" \:';R" ,'I ",\: ~ R::_-- R' R-"I -R.' 'II N

(j Figure 16·1 Diagram illustrating the inversion of NHg •

Section 15-1. The N-centered triangular metal complexes such as [NIr3(S04)6(H20)3]4- are similar.

N-N Single-Bond Energy The N-N single bond is quite weak. If we compare the single-bond energies:

HO-OH F-F Units 350 160 140 150 kJ mol-1

it is clear that there is a profound drop between C and N. This difference is prob­ ably attributable to the effects of repulsion between nonbonding lone pairs. The result is that, unlike carbon, nitrogen has little tendency to catenation.

Multiple Bonds The propensity of nitrogen, like carbon, to form jJTt-jJTt multiple bonds is a fea­ ture that distinguishes it from phosphorus and the other Group VB(l5) ele­ ments. Thus nitrogen as the element is dinitrogen (N2), with a very high bond strength and a short internuclear distance (1.094 A), whereas phosphorus forms p4 molecules or infinite layer structures in which there are only single bonds (Section 8-5). Where a nitrogen atom forms one single and one double bond, nonlinear molecules result, as shown in Structures 16-1 to 16-IV.

R1 R2 "'/ C II N. /. HO 16-1

Each nitrogen atom in these structures is of the AB2E type, uses Sp2 hybrid or­ bitals, and forms a 1t bond using the unhybridized 2p orbital. 16-3 Nitrides 401

In the oxo anions NO; (AB2E) and N0:3 (AB3), there are multiple bonds that may be formulated in either resonance or MO terms, as discussed in Chapter 3.

16-2 Occurrence and Properties of the Element

Nitrogen occurs in nature mainly as dinitrogen, N2 (bp 77.3 K), which comprises 78% by volume of the earth's atmosphere. The isotopes 14N and 15N have an ab­ solute ratio 14N/15N = 272.0. Compounds enriched in 15N are used in tracer stud­ ies. The heat of dissociation of N2 is extremely large.

1 120 t:.H= 944.7 kJ mol- K 25 •C = 10- (16-2.1)

The great strength of the N-N bond is principally responsible for the chemical inertness of N2 and for the fact that most simple nitrogen compounds are en­ dothermic even though they may contain strong bonds. Dinitrogen is notably unreactive in comparison with isoelectronic, triply bonded systems such as X-C-C-X, :C=O:, X-C-N:, and X-N=C:. Both -C C- and -C-N groups can act as donors by using their n; electrons, whereas N2 does not. It can, however, form complexes similar to those formed by CO, although to a much more limited extent, in which there are Mf-N= : and Mf-C=O: configura­ tions (Chapter 28). Nitrogen is obtained by liquefaction and fractionation of air. It usually con­ tains some argon and, depending on the quality, upwards of about 30 ppm of oxygen. Spectroscopically pure N2 is made by thermal decomposition of sodium or barium azide.

(16-2.2)

The only reactions of 2 at room temperature are with metallic Li to give Li3N, with certain transition metal complexes, and with nitrogen fixing bacteria. These nitrogen fixing bacteria are either free living or symbiotic on the root nod­ ules of clover, peas, beans, and the like. The mechanism by which these bacteria fix N2 is unknown. At elevated temperatures nitrogen becomes more reactive, especially when catalyzed. Typical reactions are

3 2 N2(g) + 3 H 2(g) = 2 H3(g) K 25 •C = 10 atm- (16-2.3) 31 N2(g) + 02(g) = 2 NO(g) K 25 •C = 5 X 10- (16-2.4) N2(g) + 3 Mg(s) = Mg3N2(s) (16-2.5) N2(g) + CaC2(s) = C(s) + CaNCN(s) (16-2.6)

16-3 Nitrides Nitrides of electropositive metals have structures with discrete nitrogen atoms and can be regarded as ionic, for example, (Ca2+)3(N3-h and (Li+)3N3-. Their ready hydrolysis to ammonia and the metal hydroxides is consistent with this. 402 Chapter 16 / Nitrogen

Such nitrides are prepared by direct interaction or by loss of ammonia from amides on heating, for example,

(16-3.1 )

Transition metal nitrides are often nonstoichiometric and have nitrogen atoms in the interstices of close-packed arrays of metal atoms. Like the similar carbides or borides they are hard, chemically inert, high melting, and electrically conducting.

There are numerous covalent nitrides (BN, S4N4' PSN 5, etc.), and their prop­ erties vary greatly depending on the element with which nitrogen is combined. These are, therefore, discussed more fully under the appropriate element.

16-4 Nitrogen Hydrides Ammonia

Ammonia (NHs) is formed by the action of a base on an ammonium salt.

(16-4.1)

Industrially, ammonia is made by the Haber process in which the reaction

!1H = -46 kJ mol-1 (16-4.2) 2 • K 25 C = lOS atm-

is carried out at 400-500 °c and pressures of 102-10s atm in the presence of a catalyst. Although the equilibrium is most favorable at low temperature, even with the best catalysts, elevated temperatures are required to obtain a satisfactory rate. The best catalyst is a-iron containing some oxide to widen the lattice and enlarge the active interface. Ammonia is a colorless, pungent gas (bp -33.35 0C). The liquid has a large heat of evaporation (1.37 kJ g-l at the boiling point) and can be handled in or­ dinary laboratory equipment. Liquid NHs resembles water in its physical behav­ ior, being highly associated via strong hydrogen bonding. Its dielectric constant

(-22 at -34°C; cf. 81 for H 20 at 25°C) is sufficiently high to make it a fair ion­ izing solvent. Its self-ionization has been discussed previously (Section 7-3).

Liquid NHs has lower reactivity than H 20 toward electropositive metals and dissolves many of them (Section 10-3). Because NHs(f) has a much lower dielectric constant than water, it is a bet­ ter solvent for organic compounds but generally a poorer one for ionic inor­ ganic compounds. Exceptions occur when complexing by NHs is superior to that by water. Thus AgI is exceedingly insoluble in water but very soluble in NHs. Primary solvation numbers of cations in NHs appear similar to those in H 20, for 2 s example, 5.0 ± 0.2 and 6.0 ± 0.5 for Mg + and Al +, respectively. Ammonia burns in air:

228 K25 .C = 10 (16-4.3) 16-4 Nitrogen Hydrides 403

Reaction 16-4.3 is thermodynamically favored under normal conditions. However, at 750-900 °C, in the presence of a platinum or a platinum-rhodium catalyst, reaction ofammonia with oxygen can be made to give NO instead ofN2, as in Eq. 16-4.4:

(16-4.4)

thus affording a useful synthesis of NO. The latter reacts with an excess of O 2 to produce N02, and the mixed oxides can be absorbed in water to form nitric acid.

2 NO + O 2 ------'? 2 N02 (16-4.5)

3 O 2 + H 20 ------'? 2 HN03 + NO and so on (16-4.6)

Thus the sequence in industrial utilization of atmospheric nitrogen is as follows:

N H 2 NH O2 ) NO (16-4.7) 2 Haber ) 3 Ostwald process process

Ammonia is extremely soluble in water. Although aqueous solutions are gen­ erally referred to as solutions of the weak base NH40H, called ammonium hy­ droxide, undissociated NH4 0H probably does not exist. The solutions are best de­ scribed as NH3 (aq), with the equilibrium written as

K = [NH:][OH-] (16-4.8) 25°C [NHg] 5 = 1.77xI0- (pKb =4.75)

Ammonium Salts

Stable crystalline salts of the tetrahedral H~ ion are mostly water soluble. Ammonium salts generally resemble those of potassium and rubidium in solubil­ ity and structure, since the three ions are of comparable (Pauling) radii: NH/ = 1.48 A, K+ = 1.33 A, Rb+ = 1.48 A. Salts of strong acids are fully ionized, and the solutions are slightly acidic.

NH4 Cl = NH; + Cl­ K""oo (16-4.9)

10 NH; + H 20 = NH3 + H 30+ K25 0C = 5.5 X 10- (16-4.10)

Thus, a 1 M solution will have a pH of about 4.7. The constant for the second re­ action is sometimes called the hydrolysis constant; however, it may equally well be considered as the acidity constant of the cationic acid NH4 +, and the system regarded as an acid-base system in the following sense:

NH: + H 20 = H 30+ + NH3 (aq) (16-4.11) Acid Base Acid Base

Many ammonium salts volatilize with dissociation around 300°C, for exam­ ple, 404 Chapter 16 I Nitrogen

!1H = 177 kJ mol-I 16 K 25 0C = 10- (16-4.12) !1H = 171 kJ mol-I (16-4.13)

Salts that contain oxidizing anions may decompose when heated, with oxidation of the ammonia to N20 or N2, or both. For example,

(NH4)2Cr207(S) = N2(g) + 4 H 20(g) + Cr20S(s) !1H = -315 !0 mol-I (16-4.14) I NH4NOsU;') = N20(g) + 2 H20(g) !1H= -23 kJ mol- (16-4.15)

Hydrazine

Hydrazine (N2H4) may be thought of as derived from ammonia by replacement of a hydrogen atom by the NH2 group. It is a bifunctional base,

7 N2H4(aq) + H20 = N2H; + OH­ K25 0C = 8.5 X 10- (16-4.16) N2H~+ 15 N2H;(aq) + H20 = + OH- K25 0C = 8.9 X 10- (16-4.17)

and two series of hydrazinium salts are obtainable. Those of N2H; are stable in water, while those of N2i-I~+ are extensively hydrolyzed. Salts of N2H~+ can be ob­ tained by crystallization from aqueous solution containing a large excess of the acid, since they are usually less soluble than the monoacid salts. Anhydrous 2H4 is a fuming colorless liquid (bp 114°C). It is surprisingly stable in view of its endothermic nature (!1H;= 50 kJ mol-I). It burns in air with considerable evolution of heat.

!1Ho = -622 kJ mol-I (16-4.18)

Aqueous hydrazine is a powerful reducing agent in basic solution, normally being oxidized to nitrogen. Hydrazine is made by the interaction of aqueous am­ monia with sodium hypochlorite.

NHs + NaOCI ~ NaOH + NH2Cl (Fast) (16-4.19) NHs + NH2Cl + NaOH ~ N2H4 + NaCI + H20 (16-4.20)

However, there is a competing reaction that is rather fast once some hydrazine has been formed.

(16-4.21)

To obtain appreciable yields, it is necessary to add gelatine. This sequesters heavy metal ions that catalyze the parasitic reaction; even the part per million or so of Cu2+ in ordinary water will almost completely prevent the formation of hy­ drazine if no gelatine is used. Since simple sequestering agents such as EDTA are not as beneficial as gelatine, the latter is assumed to have a catalytic effect as well. 16-5 Nitrogen Oxides 405

Hydroxylamine

Hydroxylamine (NH20H) is a weaker base than NHs:

9 K25 0C = 6.6 X 10- (16-4.22)

It is prepared by reduction of nitrates or nitrites either electrolytically or with S02, under controlled conditions. Hydroxylamine is a white unstable solid. In aqueous solution, or as its salts [NHsOHJCI or [NHsOHJ2S04' it is used as a reducing agent.

Azides Sodium azide can be obtained by the reaction

175°C ------7) 3 NaNH2 +NaN02 NaNs +3 NaOH+NHs (16-4.23)

Heavy metal azides are explosive and lead or mercury azide have been used in detonation caps. The azide ion, which is linear and symmetrical, behaves rather like a halide ion and can act as a ligand in metal complexes. The pure acid (HNs) is a dangerously explosive liquid.

16·5 Nitrogen Oxides Dinitrogen Monoxide (Nitrous Oxide)

Nitrous oxide (N20) is obtained by thermal decomposition of molten ammo­ nium nitrate.

(16-5.1)

The contaminants are NO (which can be removed by passage through fer­ rous sulfate solution), and 1-2% of nitrogen. Thermodynamically, nitrous oxide is unstable relative to N2 and atomic oxygen (llG= 105 kJ mol-I), but it is kinet­ ically stable in the absence of transition metal complexes with which it reacts by ° atom transfer, giving N2 and M=O or M-O-M bonds. Nitrous oxide has the linear structure NNO. It is relatively unreactive, being inert to the halogens, alkali metals, and ozone at room temperature. On heat­ ing, it decomposes to N2 and 02. At elevated temperatures, it will react with the alkali metals and with many organic compounds. It will oxidize some low-valent transition metal complexes and itself forms the complex, [Ru(NHs)sN20 f+. It is used as an anaesthetic.

Nitrogen Monoxide (Nitric Oxide) Nitric oxide (NO) is formed in many reactions involving reduction of nitric acid and solutions of nitrates and nitrites. For example, with 8 M nitric acid, we have:

(16-5.2) 406 Chapter 16 / Nitrogen

Reasonably pure NO is obtained by the aqueous reactions:

2Na O 2 +2 al+4H2S04~ 12 + 4 NaHS04 + 2 H 20 + 2 NO (16-5.3) 2 NaN02 + 2 FeS04 + 3 H 2S04 ~ Fe2(S04)3 + 2 NaHS04

+ 2 H 20 + 2 NO (16-5.4)

or, using molten salts,

(16-5.5)

Nitric oxide reacts rapidly with dioxygen, as in Reaction 16-5.6:

(16-5.6)

but the reaction is slow under dilute conditions. Nitric oxide apparently plays a respiratory role in controlling blood pressure. Nitric oxide is oxidized to nitric acid by strong oxidizing agents; the reaction with permanganate is quantitative and provides a method of analysis. It is re­ 2 duced to N20 by S02 and to NH20H by Cr +, in acid solution in both cases. Nitric oxide is thermodynamically unstable and at high pressures it readily decomposes in the range 30-50 ac.

(16-5.7)

The NO molecule is paramagnetic with the electron configuration

The electron in the 1t* orbital is relatively easily lost to give the nitrosonium ion (NO+), which forms many salts. Because the electron removed comes out of an antibonding orbital, the bond is stronger in NO+ than in 0; the bond length decreases by 0.09 Aand the vibration frequency rises from 1840 cm-l in 0 to 2150-2400 cm-l (depending on environment) in NO+. The ion is formed when N20 3 or N20 4 is dissolved in concentrated sulfuric acid.

N20 3 + 3 H 2S04 = 2 NO+ + 3 HSOi + H 30+ (16-5.8) N20 4 + 3 H 2S04 = NO+ + NO; + 3 HSOi + H 30+ (16-5.9)

The compound NO+HSOi, nitrosonium hydrogen sulfate, is an important in­ termediate in the lead-chamber process for manufacture of sulfuric acid. ot only does the NO+ ion react with many reducing agents, but it may be part of a reversible electrode reaction in nonaqueous solvents (e.g., CH3CN), as in Reaction 16--5.10.

(16-5.10)

Nitric oxide forms many complexes with transition metals (Chapter 28) some of which can be considered to arise from NO+. 16-5 Nitrogen Oxides 407

Nitrogen Dioxide (N02 ) and Dinitrogen Tetroxide (N 2 0 4 )

The two oxides, N02 and N20 4 , exist in a strongly temperature-dependent equi­ librium

2 N02 N 20 4 (16-5.11) Brown Colorless paramagnetic diamagnetic

th in solution and in the gas phase. In the solid state, the oxide is wholly N20 4 . the liquid phase, partial dissociation occurs; it is pale yellow at the freezing int (-11.2 0c) and contains 0.01% ofN02, which increases to 0.1% in the ep red-brown liquid at the boiling point, 21.15 °C. Dissociation is complete in e vapor above 140°C. Nitrogen dioxide has an unpaired electron. The other 'ee radical" molecules, NO and CI02 (Section 20-4), have little tendency to merize, and the difference may be that in N02 the electron is localized mainly 1 the N atom. The dimer has three isomeric forms ofwhich the most stable and )rmal form has the planar structure 02N-N02' The N- bond is rather long ,.75 A), as would be expected from its weakness. The dissociation energy of 204 is only 57 kJ mol-I. Mixtures of the two oxides are obtained by heating metal nitrates, by oxida­ on of NO, and by reduction of nitric acid and nitrates by metals and other re­ Lucing agents. The gases are highly toxic and attack metals rapidly. They react vith water as in Reaction 16-5.12.

(16-5.12)

[he nitrous acid decomposes, particularly when warmed:

(16-5.13)

Thermal decomposition of N02 takes place above 150°C according to Reaction 16-5.14:

(16-5.14)

The oxides are fairly strong oxidizing agents in aqueous solution, compara­ ble in strength to bromine.

EO = +1.07 V (16-5.15)

An equilibrium mixture of the oxides, nitrous fumes, is used in organic chemistry as a selective oxidizing agent, the oxidation proceeding through an initial hy­ drogen abstraction to give HO °according to Reaction 16-5.16.

RH + N02 = R· + HONO (16-5.16)

Liquid N20 4 can be used as a solvent and has been utilized to make anhy­ drous nitrates and nitrate complexes. Thus Cu dissolves in N20 4 to give CU(NOg)2'N204' which loses N20 4 on heating to give Cu(NOgh 408 Chapter 16 I Nitrogen

In solvents such as anhydrous HN03 , N 20 4 dissociates ionically as in Reaction 16-5.17.

(16-5.17)

Dinitrogen Trioxide Dinitrogen trioxide, 203' formally the anhydride of nitrous acid, is obtained by

interaction of stoichiometric amounts of NO and 02' or of NO and N 20 4 • It is an intensely blue liquid or a pale blue solid. The stable form has a weak N-N bond. It exists only at low temperature, and readily dissociates to give NO and

N02 as in Reaction 16-5.18.

(16-5.18)

The N 20 3 molecule has an 02N-NO structure in the gas phase and at low tem­ perature, with an extremely long (1.89 A.) N-N bond (Structure 16-V) consis­ tent with its easy dissociation.

1.89 A N~r ~O

16-V

Dinitrogen Pentoxide

This oxide (N 20 S) forms unstable colorless crystals. It is made by Reaction 16-5.19.

(16-5.19)

Dinitrogen pentoxide is the anhydride of nitric acid. In the solid state it exists as the nitronium nitrate, NO;N03".

16-6 The Nitronium Ion

Just as NO readily loses its odd electron, so does N02. The nitronium ion (NO;) is involved in the dissociation of HN03 , in solutions of nitrogen oxides in acids, . and in nitration reactions of aromatic compounds. Indeed, it was studies on ni­ tration reactions that lead to recognition of the importance of NO; as the at­ tacking species. 16-8 Nitrogen Halides 409

The nitronium ion is formed in ionizing solvents such as H 2S04 , CH3N02, or CH3C02H, by ionizations such as

2 HN03 = NO; + NOs + H 20 (16-6.1) HSO~ HN03 + H 2S04 = NO; + + H 20 (16-6.2)

The actual nitration process can then be formulated

itronium salts can be readily isolated. These salts are thermally stable but rapidly hydrolyzed. Typical preparations are

NO;CIO~ N20 S + HCl04 = + HN03 (16-6.4) HN03 + 2 S03 = NO;HS20; (16-6.5)

16·7 Nitrous Acid

Solutions of the weak acid HONO (pKa = 3.3) are made by acidifying cold solu­ tions of nitrites. The aqueous solution can be obtained free of salts by the reac­ tion

(16-7.1)

The pure liquid acid is unknown, but it can be obtained in the vapor phase. Even aqueous solutions of nitrous acid are unstable and decompose rapidly when heated.

(16-7.2)

Nitrites of the alkali metals are prepared by heating the nitrates with a re­ ducing agent, such as carbon, lead, or iron. They are very soluble in water. Nitrites are very toxic but have been used for preservation of ham and other meat products; there is evidence that they can react with proteins to give car­ cinogenic nitrosamines. The main use of nitrites is to generate nitrous acid for the synthesis of or­ ganic diazonium compounds from primary aromatic amines. Organic derivatives of the N02 group are of two types: nitrites (R-O 0) and nitro compounds (R-N02). Similar isomerism occurs in metal complexes where the N02"ligand can be coordinated to a metal either through the nitrogen atom (i.e., the nitro ligand) or through the oxygen atom (i.e., the nitrito ligand), as has already been discussed in Chapter 6.

16·8' Nitrogen Halides

Of the binary halides we have NF3' NF2Cl, NFCI2, and NCl3. There are also N2F2, N2F4 , and the halogen azides XN3 (X = F, CI, Br, I). With the exception of NF3, 41 0 Chapter 16 / Nitrogen

the halides are reactive and some of them are explosive, for example, NFCI2. Only the fluorides are important.

Nitrogen trifluoride is made by the electrolysis of H 4 F in anhydrous HF sol­ vent, a procedure that also gives small amounts of N2F2 . Electrolysis of molten

NH4F is the preferred method for synthesis ofN2F2 . Reaction 16-8.1,

(16-8.1)

conducted in a Cu-packed reactor, gives mixtures of fluorides. The predominant

product depends on conditions, especially the F2 /NH3 ratio. Nitrogen trifluoride (bp -129°C) is a very stable gas that normally is reac­

tive only at 250-300 DC, although it reacts readily with AlCl3 at 70°C, as in Reaction 16-8.2.

(16-8.2)

It is unreactive towards water and most other reagents at room temperature, and

it is thermally stable in the absence of reducing metals. The NF3 molecule is pyramidal, but unlike ammonia, has a very low dipole moment. Evidently, it is an extremely poor donor molecule, and does not form complexes.

Interaction of NF3 , F2, and a strong Lewis acid, such as BF3 , AsF5, or SbF5, gives salts of the ion NF/. Such reactions are performed at low temperature, under high pressures, with UV light, as in Reaction 16-8.3:

~ NF3 + F2 + BF3 NF;BFi (16-8.3)

Compounds of NF; are ionic, and other salts may be prepared similarly, namely, those of AsF6 and SnF~-. The perchlorate may be prepared by low-temperature (-78°C) metathesis in liquid HF, as in Reaction 16-8.4:

(16-8.4)

NF; is one of the strongest oxidizers known. The oxohalides (or the nitrosyl halides), XNO, where X = F, CI, or Br, are obtained by reaction of the halogens with NO as in Reaction 16-8.5.

~ 2 NO + X 2 2 XNO (16-8.5)

All three of the nitrosyl halides are powerful oxidants, able to attack many met­

als. All decompose on treatment with water producing HN03 , HN02 , NO, and HX.

16·9 Descriptive Summary of Reactions The chemistry of nitrogen is well organized by noting the oxidation state of ni­ trogen among reactants and products. As a partial summary, this is illustrated for the oxides of nitrogen in Figs. 16-2 and 16-3. The corresponding balanced chem­ ical equations are given in Tables 16-1 and 16-2. The student is encouraged to prepare similar diagrams for the hydrides and fluorides. 16-9 Descriptive Summary of Reactions 411

Oxidation State of itrogen

v

IV

III

II

o

-III Figure 16·2 Reactions that do not involve disproportionation of the oxides of nitrogen. The oxidation state of nitrogen is indicated on the scale at left.

Oxidation State of Nitrogen

V \0, NOg NO; bl ~ d IV N02 N 20 4 e a 1 b11 c NO+ III HN02 1d II NO 1 e N 20

Figure 16·3 Reactions that do involve disproportionation of the oxides of nitrogen. The oxidation state of nitrogen is indicated on the scale at left. 412 Chapter 16 I Nitrogen

Table 16-1 Chemical Equations for the Reactions of Fig. 16-2

(a) 2 O 2 + Og ----> N20 S + O2 (b) 2 N02 + H 20 2 ----> 2 H Og

(c) 2 N02 = N20 4

(d) N20 4 + xsCu ----> Cu(NOgh(s)

(e) 2 HN02 + 2 HI ----> 12 + 2 NO + 2 H20 2 g (f) Fe ++ H O 2 + H+ ----> Fe ++ NO + H 2 0

(g) 2 NO + O 2 ----> 2 N02

(h) 2 Cu + N02 ----> Cu20 + NO

(i) C + N02 ----> CO2 + ,J-N2

U) N02 + 2 H 2 ----> ,J-N2 + 2 H 20

(k) 2 N02 + 7 H 2 ----> 2 NHg + 4 H 20

(I) N2 + 3 H 2 ----> 2 NHg

(m) 4 NHg + 3 O 2 ----> 2 N2 + 6 H 20

(n) 4 NHg + 5 O 2 ----> 4 NO + 6 H 20 (0) 2 + O 2 ----> 2 NO

Table 16-2 Chemical Equations for the Reactions of Fig. 16-3

(a) 2 N02 + H 2 0 ----> HNOg + HN02

(b) N20 4 ----> 0+ + NO;;­

(c) N20 4 + 3 H2S04 ----> 0+ + NO; + 3 HS04 + HgO+

(d) 3 HN02 ----> HNOg + 2 NO + H 20

(e) 3 NO ----> N20 + N02

STUDY GUIDE

Study Questions A. Review 1. Give the electronic structure of the nitrogen atom and list the ways by which the octet can be completed in forming compounds of nitrogen. Give examples. 2. Draw the Lewis diagrams and explain the geometry and hybridization at each atom in N02, NO;;, 0;, O+, 0, 2' N;;, F 0, and N20. 3. Write balanced equations for the synthesis of nitric acid from Hg and 02'

4. Write equations for the action of heat on (a) NaNOg, (b) NH4 NOg, (c) CU(NOg)2'nHp, (d) N20, and (e) N20 g. 5. How is hydrazine prepared? 6. Write balanced equations for three different preparations of nitric oxide. 7. How is the nitronium ion prepared? Explain its significance in the nitration of aro­ matic hydrocarbons. 8. In acid solution we have

HN02 + H+ + e- ----> NO + H 20 EO = 1.0 V Write balanced equations for the reactions of nitrous acid with (a) 1-, (b) Fe2+, (c) C20~-. Study Guide 413

9. How can N02" and NOs be bonded to transition metal complexes? 10. Write balanced equations for the hydrolysis of (a) calcium nitride, (b) lithium ni­ tride, (c) dinitrogen pentoxide, and (d) dinitrogen trioxide. 11. Draw Lewis diagrams for the radicals NO and N02. Explain the formation of the -N bond in N20 S and in N20 4.

B. Additional Exercises

1. Use MO theory to compare the electronic structures of CO, N2, CN-, and NO+. Why does nitrogen form complexes with metals less readily than CO? 2. Why does nitrogen form only a diatomic molecule unlike phosphorus and other el­ ements of Group VB(l5)? 3. Nitrogen trichloride is an extremely dangerous explosive oil, but NFs is a stable gas that reacts only above 250°C. Explain this difference. 4. Three isomers of N20 4 are known. Draw likely structures for them. 5. Determine the oxidation numbers of the atoms in the molecules and ions found in the following: Reactions 16-6.1, 16-5.17, 16-5.14, 16-5.7, 16-5.1, 16-4.13, 16-4.4, and 16-2.4. Which of these are redox reactions? 6. With drawings, show how hybrid orbitals overlap in the formation of the a-bond

framework in each of the following molecules and ions: (a) N2 , (b) Ns , (c) N02", and (d) ClNO. 7. With drawings, show the hybrid orbitals that house lone pairs of electrons in the mol­ ecules and ions of Problem 6, in Part B. Be careful to show the geometry correctly, including the likely position of the lone electrons in the molecules and ions, and specify the type of hybrid that is used in each case. 8. With drawings, show the formation of the n-bond system in the molecules and ions of Problem 6, in part B. 9. Complete and balance the following equations: (a) Li + N2 (b) Cu + N02 (c) C + N02 (d) H 20 2+ N02 (e) Os + N02 (f) H2 + N02 (g) HI + HN02 10. Draw the Lewis diagrams for 2F2 and 2F4' each of which has a nitrogen-nitrogen linkage. Classify each nitrogen atom according to the ABxEyscheme of Chapter 3 and give the hybridization for each nitrogen. 11. Give the ABxEyclassification (Chapter 3) for each oxide mentioned in Section 16-5. 12. Use the style of Figs. 16-2 and 16-3 to diagram the conversion ofN02 into (a) HNOs

(b) N 20 4

(c) N 2 (d) NHs 13. Give diagrams for the stepwise conversion of N02 into (a) H O 2 and NO (b) N2 and NHs 14. Give the principal products on reaction of each of the following: (a) NHs + O 2 (uncatalyzed) (b) Disproportionation of NO. (c) Oxidation of copper by 02' (d) Oxidation ofN02 by ozone. 414 Chapter 16 I Nitrogen

(e) Reduction of N02 by excess hydrogen. (f) Disproportionation of HN02. (g) The Haber process for ammonia. (h) Hydrolysis ofN20 g .

(i) Hydrolysis of N20 S '

(j) Dissolution ofN20 4 in anhydrous HNOg . (k) Dimerization of N02. (I) Oxidation of ammonia by air over a Pt catalyst at 750°C.

15. Outline the synthesis of HNOg , starting from the elements.

16. Which oxide is the anhydride of HNOg ? of HN02? Explain by using equations.

C. Questions from the Literature of Inorganic Chemistry

1. Hydrolysis of dinitrogen trioxide is described in the paper by G. Y. Markovits, S. E. Schwartz, and L. Newman, Inorg. Chern., 1981, 20, 445-450. (a) Draw the Lewis diagrams, discuss the geometry, and assign an oxidation number to each atom in the substances found in Reactions (I), (2), and (3).

(b) What evidence do the authors cite for an equilibrium in which N20 g is formed from nitrous acid in acidic medium? (c) How is Eq. (15) obtained? (d) The authors report a value for ~G/[N203(aq)]. How was this number calcu­ lated? (e) Of Reactions (1), (2), (3), (9), and (13), which represent disproportionation, hydrolysis, and/or acid-base type reactions? 2. Consider the paper by K. 0. Christe, C. J. Schack, and R. D. Wilson in Inorg. Chem., 1977, 16, 849-854.

(a) What is the nature of solid SnF4 , and why is it not a good Lewis acid?

(b) What reaction takes place in liquid HF solvent between KF and SnF4 ?

(c) Draw the Lewis diagrams and predict the geometries of NF;, BF4, (SnF4 )x, SnFs,and [SnF6f-.

(d) What reaction takes place in liquid HF solvent between NF4 BF4 and SnF4 ?

(e) Why does NF4SnFs not react with a second equivalent of F4BF4 to form

(NF4 )2SnF6? 3. Some reaction chemistry of NF; is reported by K. 0. Christe, W. W. Wilson, and R. D. Wilson in Inorg. Chern., 1980, 19, 1494-1498. (a) Write balanced equations for the reactions in anhydrous HF solvent between NF; and (i) CI04, (ii) Br04, and (iii) HFi. (b) Write balanced equations for the reactions in BrFs solvent between NF; and (i) BrF4and (ii) BrF4 0-.

SUPPLEMENTARY READING

Bottomley, F. "Reactions of Nitrosyls," in Reactions of Coordinated Ligands, P. S. Braterman, Ed., Plenum, New York, 1989. Bottomley, F. and Burns, R. c., Treatise on Dinitrogen Fixation, Wiley, New York, 1979. Chatt, J. C., da C. Pina, L. M., and Richards, R. L., New Trends in Nitrogen Fixation, Academic, New York, 1980. Supplementary Reading 415

Colburn, C. B., Ed., Developments in Inorganic Nitrogen Chemistry, Vols. 1 and 2, Elsevier, Amsterdam, 1966 and 1973. Dehnicke, K. and Strahl, J., "Nitrido Complexes of the Transition Metals," Angew. Chem. !nt. Ed. Eng., 1992, 31,955-978. Emeleus, H.J., Shreeve,J. M., and Verma, R. D., "The Nitrogen Fluorides and Some Related Compounds," Adv. Inorg. Chem., 1989, 33, 139-196. Griffith, W. P., "Transition-Metal Nitrido Complexes," Coord. Chem. Rev., 1972, 8, 369-396. Jolly, W. L., The Inorganic Chemistry ofNitrogen, Benjamin, ew York, 1964. Smith, P. A. S., The Open-Chain Chemistry of Organic Nitrogen Compounds, Vols. 1 and 2, Benjamin, New York, 1966. Wright, A. N. and Winkler, C. A., Active Nitrogen, Academic, New York, 1968.