Chapter 16
NITROGEN
16-1
Introduction
The nitrogen atom (ls22s22Px2py2pz) can complete its valence shell in the fol­
lowing ways:
1. Electron gain to form the nitride ion N 3 -; this ion is found only in saltlike
nitrides of the most electropositive metals.
2. Formation of electron-pair bonds: (a) single bonds, as in NH 3 , or (b)
multiple bonds, as in :N=N:, -N=N-, or N0 2 •
3. Formation of electron-pair bonds with electron gain, as in NH 2 or NH 2 -.
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 3 E (as in pyramidal NR3 ); AB 2 E 2 (as in bent NR2); AB 3 (as in planar NO g);
AB 2 E (as in bent R2 C=N-OH); and ABE (N 2 ). There are a few stable species
in which, formally, the nitrogen valence shell is incomplete. Nitroxides,
R 2 N=Q, NO, and N0 2 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, F3 B:N(CH 3 h, and they act as ligands toward transition
metal ions as in, for example, [Co(NH 3 )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 (SiMe 3 ) 3' in
399
400
Chapter 16
/
Q'
N
I
\:';R"
",\: ~
R-"-R.'
Nitrogen
R"
,'I
R::_-R'
'II
N
(j
Figure 16·1
Diagram illustrating the inversion of NH g •
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:
350
160
HO-OH
F-F
Units
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 (N 2), with a very high bond
strength and a short internuclear distance (1.094 A), whereas phosphorus forms
p 4 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 AB 2E type, uses
bitals, and forms a 1t bond using the un hybridized 2p orbital.
Sp2
hybrid or­
16-3
401
Nitrides
In the oxo anions NO; (AB 2E) and N0:3 (AB 3), 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, N 2 (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 N 2 is extremely large.
t:.H= 944.7 kJ mol- 1
K 25 • C = 10- 120 (16-2.1)
The great strength of the N-N bond is principally responsible for the chemical
inertness of N 2 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 N 2 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 N 2 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
Li 3N, 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 N 2 is unknown.
At elevated temperatures nitrogen becomes more reactive, especially when
catalyzed. Typical reactions are
N 2(g) + 3 H 2(g)
=2
H 3(g)
N 2(g) + 02(g) = 2 NO(g)
16-3
K 25 • C = 103 atm- 2
K 25 • C = 5
X
10- 31
(16-2.3)
(16-2.4)
N 2(g) + 3 Mg(s) = Mg 3N 2(s)
(16-2.5)
N 2(g) + CaC 2(s) = C(s) + CaNCN(s)
(16-2.6)
Nitrides
Nitrides of electropositive metals have structures with discrete nitrogen atoms
and can be regarded as ionic, for example, (Ca2+)3(N 3-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' P SN 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 (NH s ) 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)
K 25 • C = lOS atm- 2
is carried out at 400-500 °c and pressures of 10 2-10 s 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 NH s 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 2 0 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 NH s has lower reactivity than H 2 0 toward electropositive metals and
dissolves many of them (Section 10-3).
Because NH s (f) has a much lower dielectric constan t 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 NH s is superior to that
by water. Thus AgI is exceedingly insoluble in water but very soluble in NH s.
Primary solvation numbers of cations in NH s appear similar to those in H 2 0, for
example, 5.0 ± 0.2 and 6.0 ± 0.5 for Mg 2 + and Al s+, respectively.
Ammonia burns in air:
K 25
.C
=
10 228 (16-4.3)
16-4
403
Nitrogen Hydrides
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 of ammonia with oxygen can be made to give NO instead ofN 2 ,
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 N0 2 , and the mixed oxides can be absorbed in water to form nitric acid.
3
2 NO + O 2
-------'?
2 N0 2
O 2 + H 20
-------'?
2 HN0 3 + NO
(16-4.5)
and so on
(16-4.6)
Thus the sequence in industrial utilization of atmospheric nitrogen is as follows:
N2
H2
Haber )
process
NH
O2
3
Ostwald
process
)
NO
(16-4.7)
Ammonia is extremely soluble in water. Although aqueous solutions are gen­
erally referred to as solutions of the weak base NH 4 0H, called ammonium hy­
droxide, undissociated NH 4 0H probably does not exist. The solutions are best de­
scribed as NH 3 (aq), with the equilibrium written as
K
=
[NH:][OH-]
[NH g ]
=
1.77xI0-5 (pK b =4.75)
25°C
(16-4.8)
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.
NH 4 Cl = NH; + Cl­
NH; + H 2 0
= NH 3 + H 3 0+
(16-4.9)
K""oo
K 25
0C
= 5.5
X
10-
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 NH 4 +, and the system
regarded as an acid-base system in the following sense:
NH:
Acid
+ H 20
Base
=
H 3 0+
+ NH 3 (aq)
Acid
Base
(16-4.11)
Many ammonium salts volatilize with dissociation around 300°C, for exam­
ple,
404
Chapter 16
I
Nitrogen
!1H = 177 kJ mol- I
K 25
0C
= 10- 16
(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 N 20 or N 2, or both. For example,
(NH4)2Cr207(S) = N 2(g) + 4 H 20(g) + Cr 20S(s)
!1H = -315
!0 mol- I
!1H= -23 kJ mol- I
NH 4NOsU;') = N 20(g) + 2 H 20(g)
(16-4.14)
(16-4.15)
Hydrazine
Hydrazine (N 2H 4) may be thought of as derived from ammonia by replacement
of a hydrogen atom by the NH 2 group. It is a bifunctional base,
N 2H 4(aq) + H 20 = N 2H; + OH­
K 25 0C = 8.5
X
10-7
(16-4.16)
N 2H;(aq) + H 20 = N2H~+ + OH-
K 25
X
10- 15
(16-4.17)
0C
= 8.9
and two series of hydrazinium salts are obtainable. Those of N 2H; 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.
NH s + NaOCI ~ NaOH + NH 2Cl
NH s + NH 2Cl + NaOH
~
N 2H 4 + NaCI + H 20
(Fast) (16-4.19)
(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
Cu 2+ 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
405
Nitrogen Oxides
Hydroxylamine
Hydroxylamine (NH 20H) is a weaker base than NH s :
K 25
0C
= 6.6
X
10-9
(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
3 NaNH 2 +NaN0 2 - - - - - - 7 ) NaN s +3 NaOH+NH s
(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
(HN s ) is a dangerously explosive liquid.
16·5
Nitrogen Oxides
Dinitrogen Monoxide (Nitrous Oxide)
Nitrous oxide (N 20) 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 N 2 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 N 2 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 N 2 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 (NH s )sN 20 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~
(16-5.3)
12 + 4 NaHS0 4 + 2 H 20 + 2 NO
2 NaN0 2 + 2 FeS0 4 + 3 H 2S0 4 ~ Fe 2(S04)3 + 2 NaHS0 4
+ 2 H 2 0 + 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
duced to N 20 by S02 and to NH 20H by Cr 2+, 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 A and 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 N 20 3 or N 20 4 is dissolved in concentrated sulfuric
acid.
N 20 3 + 3 H 2S0 4 = 2 NO+ + 3 HSOi + H 30+
(16-5.8)
N 20 4 + 3 H 2S0 4 = 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., CH 3CN), 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
407
Nitrogen Oxides
Nitrogen Dioxide (N0 2 ) and Dinitrogen Tetroxide (N 2 0 4 )
The two oxides, N0 2 and N 20
librium
4,
exist in a strongly temperature-dependent equi­
2 N0 2
Brown
paramagnetic
N 20
(16-5.11)
4
Colorless
diamagnetic
th in solution and in the gas phase. In the solid state, the oxide is wholly N 20 4 .
the liquid phase, partial dissociation occurs; it is pale yellow at the freezing
int (-11.2 0c) and contains 0.01% ofN0 2, 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 CI0 2 (Section 20-4), have little tendency to
merize, and the difference may be that in N0 2 the electron is localized mainly
1 the N atom. The dimer has three isomeric forms of which 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 N0 2 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 + N0 2 = R· + HONO
(16-5.16)
Liquid N 20 4 can be used as a solvent and has been utilized to make anhy­
drous nitrates and nitrate complexes. Thus Cu dissolves in N 20 4 to give
CU(NOg)2'N204' which loses N 20 4 on heating to give Cu(NOgh
408
Chapter 16
I
Nitrogen
In solvents such as anhydrous HN0 3 , N 20
Reaction 16-5.17.
4
dissociates ionically as in
(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
N0 2 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 2 0
16-5.19.
S)
forms unstable colorless crystals. It is made by Reaction
(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 N0 2. The nitronium ion (NO;)
is involved in the dissociation of HN0 3 , 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
409
Nitrogen Halides
The nitronium ion is formed in ionizing solvents such as H 2S0 4 , CH 3N0 2,
or CH 3C0 2H, by ionizations such as
2 HN0 3 = NO; + NOs + H 20
HN0 3 + H 2S0 4 = NO; +
HSO~
+ H 20
(16-6.1)
(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
N 20 S + HCl0 4 =
NO;CIO~
+ HN0 3
HN0 3 + 2 S03 = NO;HS 20;
16·7
(16-6.4)
(16-6.5)
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 N0 2 group are of two types: nitrites (R-O 0) and nitro compounds
(R-N0 2). 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, NFCI 2, and NCl 3. There are also N 2F2,
N 2F4 , and the halogen azides XN 3 (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, NFCI 2 .
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 N 2 F2 . Electrolysis of molten
NH 4 F is the preferred method for synthesis ofN 2 F2 . 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 /NH 3 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 AlCl 3 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 NF 3 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 , F 2 , and a strong Lewis acid, such as BF3 , AsF 5 , 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:
NF 3 + F2 + BF3 ~ NF;BFi
(16-8.3)
Compounds of NF; are ionic, and other salts may be prepared similarly, namely,
those of AsF 6 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 HN0 3 , HN0 2 , 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
411
Descriptive Summary of Reactions
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
\0,
V
N0 2
IV
e
III
1
1
1
a
HN0 2
NO g
bl
d
N 20
11
b
NO;
~
4
c
NO+
d
II
NO
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
Table 16-1
Chemical Equations for the Reactions of Fig. 16-2
Nitrogen
(a) 2 O 2 + Og ----> N 20 S + O 2
(b) 2 N0 2 + H 20 2 ----> 2 H Og
(c) 2 N0 2 = N 2 0 4
(d) N 2 0 4 + xsCu ----> Cu(NOgh(s)
(e)
(f)
2 HN0 2 + 2 HI ----> 12 + 2 NO + 2 H 20
g
Fe 2 + + H O 2 + H+ ----> Fe + + NO + H 2 0
(g) 2 NO + O 2 ----> 2 N0 2
(h) 2 Cu + N0 2 ----> Cu 20 + NO
(i) C + N0 2 ----> CO 2 + ,J-N 2
U)
(k)
N0 2 + 2 H 2 ----> ,J-N 2 + 2 H 2 0
2 N0 2 + 7 H 2 ----> 2 NHg + 4 H 2 0
(I) N 2 + 3 H 2 ----> 2 NHg
(m) 4 NHg + 3 O 2 ----> 2 N 2 + 6 H 2 0
(n) 4 NHg + 5 O 2 ----> 4 NO + 6 H 2 0
(0)
2 + O 2 ----> 2 NO
Table 16-2
Chemical Equations for the Reactions of Fig. 16-3
(a) 2 N0 2 + H 2 0 ----> HNOg + HN0 2
(b) N 2 0 4 ----> 0+ + NO;;­
(c) N 20 4 + 3 H 2S0 4 ----> 0+ + NO; + 3 HS0 4 + HgO+
(d) 3 HN0 2 ----> HNOg + 2 NO + H 2 0
(e) 3 NO ----> N 2 0 + N0 2
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 N0 2, NO;;,
0;,
O+,
0,
2'
N;;, F 0, and N 2 0.
3. Write balanced equations for the synthesis of nitric acid from H g and 02'
4. Write equations for the action of heat on (a) NaNOg, (b) NH 4 NOg,
(c) CU(NO g )2'nHp, (d) N 20, and (e) N 20 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
HN0 2 + H+ + e- ----> NO + H 2 0
EO = 1.0 V
Write balanced equations for the reactions of nitrous acid with (a) 1-, (b) Fe 2+, (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 N0 2. Explain the formation of the
-N bond in N 20 S and in N 20 4.
B. Additional Exercises
1. Use MO theory to compare the electronic structures of CO, N 2, 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 N 20 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) N 2 , (b) N s , (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 + N 2
(b) Cu + N0 2
(d) H 20 2 + N0 2
(c) C + N0 2
(f) H 2 + N0 2
(e) Os + N0 2
(g) HI + HN0 2
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 ofN0 2 into
(a) HNO s
(b) N 2 0
4
(c) N 2
(d) NH s
13. Give diagrams for the stepwise conversion of N0 2 into
(a) H O 2 and NO
(b) N 2 and NH s
14. Give the principal products on reaction of each of the following:
(a) NH s + O 2 (uncatalyzed)
(b) Disproportionation of NO.
(c) Oxidation of copper by 02'
(d) Oxidation ofN0 2 by ozone.
414
I
Chapter 16
Nitrogen
(e) Reduction of N0 2 by excess hydrogen.
(f) Disproportionation of HN0 2.
The Haber process for ammonia.
Hydrolysis ofN 20 g .
Hydrolysis of N 20 S '
Dissolution ofN 20 4 in anhydrous HNO g .
Dimerization of N0 2.
(I) Oxidation of ammonia by air over a Pt catalyst at 750°C.
15. Outline the synthesis of HNO g , starting from the elements.
16. Which oxide is the anhydride of HNO g ? of HN0 2? Explain by using equations.
(g)
(h)
(i)
(j)
(k)
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 N 20 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,
SnF and [SnF6f-.
(d) What reaction takes place in liquid HF solvent between NF4 BF4 and SnF4 ?
(e) Why does NF4 SnFs not react with a second equivalent of F4 BF4 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) CI0 4, (ii) Br0 4, and (iii) HFi.
(b) Write balanced equations for the reactions in BrFs solvent between NF; and (i)
BrF4 and (ii) BrF4 0-.
s,
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.
415
Supplementary Reading
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 of Nitrogen, 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.