Kinetics of oxidation of nitrogen compounds by cerium(IV)

Article abstract of DOI:10.1039/a904080c

The kinetics of oxidation by Ce^IV in sulfuric acid of hydrazine, hydroxylamine, nitrous acid, hydrazoic acid and of an intermediate involved in the nitrite/hydrazine reaction has been examined under a uniform set of conditions. Reaction proceeds through the free base form of the nitrogen substrate, probably by an inner-sphere mechanism, with a reactivity sequence N₂H₄ > NH₂OH > N₃^- > NO₂^-. An intermediate in the hydrazine/nitrous acid reaction, NH₂N=NOH, is also oxidised by Ce^IV in a Ce^IV/HNO₂/N₂H₅⁺ system. For the Ce^IV/HNO/HN₃ system there is a much larger consumption of oxidant than can be accounted for by the separate oxidations of nitrite and azide. An additional pathway is proposed, probably involving NNN-O-N=O, formed by combination of azide and nitrogen dioxide radicals. The Royal Society of Chemistry 1999.

Full text of DOI:10.1039/a904080c

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Kinetics of oxidation of nitrogen compounds by cerium(IV)
Anne M. M. Doherty,a Mark D. Radcliffeb and Geoffrey Stedman*b
a Science Department, University College of Worcester, Henwick Grove, Worcester,
UK WR2 6AJ
b Chemistry Department, University of Wales Swansea, Singleton Park, Swansea, UK SA2 8PP
Received 21st May 1999, Accepted 23rd July 1999
The kinetics of oxidation by CeIV in sulfuric acid of hydrazine, hydroxylamine, nitrous acid, hydrazoic acid and
of an intermediate involved in the nitrite/hydrazine reaction has been examined under a uniform set of conditions.
Reaction proceeds through the free base form of the nitrogen substrate, probably by an inner-sphere mechanism,
Ϫ
with a reactivity sequence N2H4 > NH2OH > N3Ϫ > NO2 . An intermediate in the hydrazine/nitrous acid reaction,
NH N᎐NOH, is also oxidised by CeIV in a CeIV/HNO2/N2H5ϩ system. For the CeIV/HNO/HN3 system there is a
᎐
2
much larger consumption of oxidant than can be accounted for by the separate oxidations of nitrite and azide. An
additional pathway is proposed, probably involving NNN–O–N᎐O, formed by combination of azide and nitrogen
᎐
dioxide radicals.
The present study began as an attempt to detect intermediates
in various inorganic N-nitrosations by reaction with CeIV using
a kinetic approach. In the analysis of an aqueous solution of
NaNO2 ϩ NaN3 in NaOH the nitrite ion was measured spec-
trophotometrically at 355 nm, and it was hoped to determine
the total nitrite ϩ azide volumetrically by reaction with an
excess of cerium() ion in sulfuric acid, and back titration of
the residual oxidant by standardised iron() ion. As nitrite and
azide react together in acid solution, eqns. (1) and (2), it was
Kinetics
Fast reactions were followed on a Canterbury SF 3A stopped-
flow instrument with a data collection system. Measurements
were made at 418 nm, calibrating with a holmium filter.
Titrimetric analyses for CeIV
In these experiments a 5 cm3 aliquot of CeIV in sulfuric acid in a
25 cm3 conical flask was vigorously stirred using a magnetic
follower. An equal volume of NaNO2 ϩ NaN3 solution was
pipetted into the stirred solution, and there was a rapid
evolution of gas. A clean dry 1 cm spectrophotometer cell was
quickly filled from the solution in the conical flask and the
absorbance measured at a pre-determined wavelength in a
UNICAM 8625 spectrometer. It was essential to make this
measurement as quickly as possible to avoid errors due to
bubbles of liberated gas sticking to the cell wall in the light
beam. The solution was then tipped back into the conical flask
which was replaced on the magnetic stirrer. After a short inter-
val, typically 30–60 s, the cell was refilled and the absorbance
remeasured. This was repeated over a period of minutes until
no further decrease in absorbance was observed. Finally, the
solution was titrated with standardised ammonium iron()
sulfate, using ferroin as an indicator. In experiments where a
nitrate analysis was planned the excess of CeIV was titrated to
a colourless end-point with standard sodium oxalate solution.
These experiments were carried out at ambient temperatures,
21–22 ЊC.
Hϩ ϩ HNO2 ϩ N3Ϫ → N3NO ϩ H2O
N3NO → N2 ϩ N2O
(1)
(2)
necessary to check whether losses by such a reaction, which
should lead to a reduced consumption of CeIV, were significant.
When standard solutions of NaNO2 ϩ NaN3 were analysed it
was found that the consumption of cerium() ion was actually
much greater than that calculated from the known initial
concentrations. This raised the possibility that CeIV was react-
ing with an intermediate, nitrosyl azide, in the nitrite/azide
reaction.
The classic work of Clusius and Effenberger1 using 15N
tracer showed that the nitrite/azide reaction probably pro-
ceeded via nitrosyl azide, a result consistent with the general
pattern of electrophilic reactions of nitrous acid2 which com-
monly involve the formation of nitrosyl compounds such as
ONCl, ONBr, ONSCN, etc. The decomposition reaction (2) is
too fast for it to be observed directly by stopped-flow, but it was
hoped to trap it by reaction with CeIV. The work was extended
to see if intermediates in other inorganic N-nitrosation
reactions could be detected. One advantage of the kinetic
approach was that it provided relative reactivities of a number
of compounds towards CeIV under a standard set of condi-
tions. There have been previous studies of the kinetics of the
cerium(IV) oxidation of inorganic nitrogen compounds, but
reaction conditions have varied considerably.
Results
Kinetic studies
Stopped-flow spectrophotometry was used to study the kinetics
of reaction of CeIV with HNO2, HN3, N2H5 , NH3OHϩ and
ϩ
NH2SO3H by following the decrease in absorbance, A, due to
CeIV at 418 nm. All reactions were followed with one reagent in
large excess over the other and gave excellent first order kinetics
over at least three half-lives with plots of ln (A Ϫ A∞) versus
time, yielding rate constants k1/sϪ1. Average standard deviations
for repeat runs were 0.7%. All reactions were run in 0.725 mol
dmϪ3 sulfuric acid at 25 ЊC. From the results of Robertson and
Dunford3 this corresponds to [Hϩ] = 0.929 mol dmϪ3 with a
formal ionic strength of 1.13 mol dmϪ3.
Experimental
Materials
All chemicals used were AnalaR materials (BDH) with the
exception of N2H5HSO4 which was an Aldrich product. They
were used without further purification.
J. Chem. Soc., Dalton Trans., 1999, 3311–3316
This journal is © The Royal Society of Chemistry 1999
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Table 1 Kinetics of reaction of CeIV with HNO2, HN3 and NH3OHϩ
in 0.725 mol dmϪ3 sulfuric acid at 25 ЊC
have studied the reaction in a large excess of hydrazine and
again find k1 directly proportional to [N2H5ϩ] yielding a second
order rate constant k5 = 13.4 0.13 dm3 molϪ1 sϪ1, which com-
pares with a value5 of 9.6 dm3 molϪ1 sϪ1 in 1 mol dmϪ3 per-
chloric acid. These are stoichiometric rate constants based
upon the total concentration of hydrazine species. If there is
102[HNO2]/mol dmϪ3
k1/sϪ1
5
18.2
5
4
15.1
4
88.7
4
3
11.4
3
65.3
3
2
7.36
2
1
3.65
1
102[HN3]/mol dmϪ3
k1/sϪ1
115
43.3
20.9
2ϩ
102[N2H5ϩ]/mol dmϪ3
102k1/sϪ1
5
67.5
2
27.1
1
13.9
significant protonation to N2H6 then these values need to be
corrected. Stanbury6 has reviewed literature data and suggested
52.8
40.5
2ϩ
a pK1 for N2H6 of Ϫ0.5 which converts our k5 from 13.4
ϩ
into 17.4 dm3 molϪ1 sϪ1 (for N2H5 as the hydrazine species).
Margerum and Jia7 have recently measured pK1 and find a
value of 0.23 at 1 mol dmϪ3 ionic strength which gives a cor-
rected k5 of 35.3 dm3 molϪ1 sϪ1
.
Reaction with NH2SO3H. There was no detectable reaction
on the stopped-flow timescale.
Acidity dependence of the above reactions
As both oxidant and reductant can exist in conjugate acid–base
forms it was necessary to know the acidity dependence of k1.
In sulfuric acid CeIV exists as a mixture of Ce4ϩ, [Ce(OH)]3ϩ
,
[Ce(SO4)]2ϩ, Ce(SO4)2 and [Ce(SO4)3]2Ϫ. At sufficiently high
concentrations of CeIV dimeric species exist. Changes in the
sulfuric acid concentration may thus vary the proportions of
different complexes present as well as varying [Hϩ]. Further-
more, the formal ionic strength in our system is outside the
range for which activity coefficient corrections can be calcu-
lated. Hanna et al.8 suggested a method of preparing a solution
of HClO4 ϩ NaClO4 ϩ Na2SO4 of a given ionic strength, in
which [Hϩ] can be varied but [SO42Ϫ] is kept constant. This
should keep the relative concentrations of the various sulfate
complexes constant. We have used Hanna’s method to com-
pare rates at [Hϩ] = 0.7 and 0.2 mol dmϪ3. For a 3.5 fold
decrease in [Hϩ] k2 increased by a factor of 3.1, k3 by 3.0, k4
by 3.1 and k5 by 3.4. If a correction for double protonation
of hydrazine is made Stanbury’s value for the pKa leads to a
corrected factor of 2.9 for k5 while Margerum’s value reduces it
to 2.1.
Fig. 1 First order plot for reaction of CeIV with an excess of HNO2 at
25 ЊC, 104[CeIV] = 5, 102[HNO2] = 5, [H2SO4] = 0.725 mol dmϪ3
.
Reaction with HNO2. This is a well established reaction for
the volumetric estimation of nitrites; the stoichiometry is shown
in eqn. (3). Reaction with a large excess of nitrous acid gave
2Ce4ϩ ϩ HNO2 ϩ H2O → 2Ce3ϩ ϩ HNO3 ϩ 2Hϩ (3)
values of k1 directly proportional to [HNO2] yielding rate
law (4) with k2 = 370 5.5 dm3 molϪ1 Ϫ1. Data are set out in
s
Ϫd[CeIV]/dt = k2[CeIV][HNO2]
(4)
Table 1 and a typical run is shown in Fig. 1. Experiments with a
large excess of CeIV gave k2 = 386 12 dm3 molϪ1 sϪ1
Kinetics of reaction in a mixed CeIV/HNO2/N2H5؉ system
.
The reaction between nitrous acid and hydrazine follows kinetic
law (8), an initial N-nitrosation to form NH2NHNO (or a
Reaction with HN3. This is also a reaction used as a standard
volumetric procedure with the 1:1 stoichiometry of eqn. (5).
Ϫd[HNO2]/dt = k6[Hϩ][HNO2][N2H5
(8)
ϩ
]
Ce4ϩ ϩ HN3 → Ce3ϩ ϩ 1.5N2 ϩ Hϩ
(5)
conjugate acid thereof), followed by a rapid tautomerisation to
The kinetics was again studied with a large excess of azide, and
k1 found to be directly proportional to [HN3] (Table 1) giving a
second order rate constant k3 = 2280 39 dm3 molϪ1 sϪ1. When
form an intermediate NH N᎐NOH that can be detected spec-
᎐
2
trophotometrically.9 To investigate whether this intermediate
would react with CeIV one reservoir of the stopped-flow appar-
atus contained NaNO2 ϩ N2H4 while the other contained CeIV
in sulfuric acid. Monitoring the nitrite ion peak at 355 nm for
the reservoir solution showed no change in absorbance over a
period of hours confirming the stability in the absence of acid.
Conditions were chosen so that the known rate of (8) was much
faster than the rate of the CeIV/HNO2 reaction so that on rapid
mixing the nitrite/hydrazine mixture was acidified and the
an excess of CeIV was used, k3 = 2211 29 dm3 molϪ1 sϪ1
.
Reaction with NH3OH؉. This has been studied in detail by
Waters and Wilson4 who observed second order kinetics, k4,
first order with respect to each reactant concentration. For an
excess of hydroxylamine the stoichiometry is shown in eqn. (6).
Ce4ϩ ϩ NH3OHϩ → Ce3ϩ ϩ 0.5N2 ϩ H2O ϩ 2Hϩ (6)
intermediate NH N᎐NOH formed almost quantitatively. Thus,
᎐
2
for 0.725 mol dmϪ3 sulfuric acid and [N2H5ϩ]0 = 0.035 mol dmϪ3
at 25 ЊC the rate of reaction (8) is Ϫd[HNO2]/dt = 26.7[HNO2]
whereas for reaction with 5 × 10Ϫ4 mol dmϪ3 CeIV the rate is
Under our standard conditions with an excess of NH3OHϩ,
and assuming the kinetic form found by Waters and Wilson, we
found k4 = 111 1.7 dm3 molϪ1 sϪ1
.
Ϫd[HNO2]/dt = 0.092[HNO2] mol dmϪ3 Ϫ1. Thus nearly all of
s
the nitrous acid reacts with the hydrazine and almost none with
CeIV. On mixing there was a rapid decrease in [CeIV], much
faster than the known rate of the CeIV/hydrazine reaction as
shown in Fig. 2. The continuous line represents the predicted
rate of consumption of CeIV if the only processes are the CeIV/
hydrazine, CeIV/nitrous acid and nitrous acid/hydrazine reac-
tions. From the half-lifetimes for the decrease in absorbances an
Reaction with N2H5؉. Morrow and Sheeres5 have shown that
for the cerium() oxidation of hydrazine in perchloric acid the
stoichiometry is dependent upon the oxidant:reductant ratio.
For an excess of hydrazine the reaction is shown in eqn. (7). We
Ce4ϩ ϩ N2H5ϩ → Ce3ϩ ϩ NH4ϩ ϩ 0.5N2 ϩ Hϩ (7)
3312
J. Chem. Soc., Dalton Trans., 1999, 3311–3316

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Table 2 Stoichiometric data on the reaction of excess of CeIV with sodium nitrite ϩ sodium azide
[CeIV]0/
[NO2Ϫ]0/
[N3Ϫ]0/
[CeIV]∞/
Run
mol dmϪ3
mol dmϪ3
mol dmϪ3
mol dmϪ3
R
k/sϪ1
1
2
3
6
7
8
9
10
12
13
18
19
20
21
22
23
24
25
26
27
28
34
36
0.125
0.125
0.083
0.0625
0.125
0.125
0.125
0.125
0.125
0.125
0.0125
0.0166
0.0227
0.0125
0.125
0.083
0.0125
0.125
0.125
0.125
0.125
0.005
0.00469
0.025
0.025
0.033
0.000
0.025
0.025
0.000
0.050
0.0005
0.0007
0.01613
0.0135
0.0008
0.0626
0.0841
0.0850
0.0418
0.0843
0.000
0.00862
0.004218
0.00638
0.0596
0.000
0.0474
0.0505
0.0554
0.0692
0.0697
0.00227
0.00136
1.66
1.66
2.02
0.98
1.66
1.66
1.64
1.60
1.66
1.63
>1.66
1.61
2.04
1.63
1.64
>1.56
1.55
1.32
1.25
1.49
1.47
1.82
1.77
0.016
0.025
0.025
0.0125
0.0083
0.0083
0.0166
0.0083
0.0025
0.0016
0.0091
0.00125
0.01
0.0125
0.0083
0.0083
0.0166
0.0083
0.0025
0.0016
0.00
0.00125
0.02
0.0266
0.01
0.00625
0.00625
0.025
0.025
0.0005
0.000625
0.018
0.018
0.017
0.013
0.015
0.014
0.013
0.0133
0.02
0.025
0.012
0.013
0.015
0.025
0.00625
0.00625
0.0005
0.000625
Fig. 2 Reaction of CeIV with N2H5ϩ ϩ HNO2 at 25 ЊC, 104[CeIV] = 5,
103[N2H5ϩ] = 35, 103[HNO2] = 5, [H2SO4] = 0.725 mol dmϪ3
Fig. 3 Reaction of CeIV with HN3 ϩ HNO2 at 25 ЊC, 103[CeIV] = 5,
104[HN3] = 2, 104[HNO2] = 3, [H2SO4] = 0.725 mol dmϪ3. x, Experi-
mental points; full line, computed model.
.
apparent order of reaction with respect to [CeIV] can be calcu-
lated and was found to be 0.88 for [N2H5ϩ]0 = 0.035 decreasing
to 0.81 for [N2H5ϩ]0 = 0.0075 mol dmϪ3. When [HNO2]0 was
varied it was found that Ϫ(d[CeIV]/dt)0 was proportional to
[HNO2]0.
Reaction of the CeIV/HNO2/NH3OH؉ and the CeIV/HNO2/
NH2SO3H systems
These experiments were unsuccessful because the reaction of
hydroxylamine with nitrous acid is relatively slow11 and we were
unable to obtain any evidence for reaction of CeIV with any of
the postulated intermediates in the HNO2/NH3OHϩ reaction.
The nitrite reacted preferentially with CeIV. Sulfamic acid is,
like hydrazine, an excellent trap for nitrous acid12,13 and it was
easy to set up conditions where the nitrite reacted almost 100%
with NH2SO3H. However, in the three reactant system there
was no consumption of CeIV; either the expected intermediates
Kinetics of the CeIV/HNO2/HN3 systems
The cerium() consumption experiments that led to the present
work involved relatively high concentrations of reagents, e.g.
[CeIV] 0.125, [NaNO2] 0.025 and [NaN3] 0.025 mol dmϪ3. At
these concentrations the rates of reaction were too fast to be
studied by our stopped-flow. Thus for the CeIV/HN3 reaction
t1/2 ≈ 2.5 ms, considerably faster than our mixing time. Fur-
thermore, vigorous gas evolution made it impossible to get
good absorbance vs. time traces. Reasonable curves could be
obtained for [CeIV] = 0.005 mol dmϪ3 with correspondingly
lower concentrations of azide and nitrite and a typical trace is
shown in Fig. 3. There was a rapid initial decrease in absorb-
ance followed by a very much slower decay. The experimental
points, shown as crosses, are compared with the full line cal-
culated from the known rates10 of reactions (1), (3) and (5).
These results contrast with those from experiments in which
nitrite and azide were in large excess over cerium where the
curve of [CeIV] against time was in close agreement with that
calculated.
ONؒNHSO H or HON᎐NSO H were resistant to oxidation or
᎐
3
3
they decomposed too rapidly.
Stoichiometric experiments on the CeIV/HNO2/HN3 system
Most of this work consisted of determining the residual [CeIV]
when a mixture of NaNO2 ϩ NaN3 was treated with an excess
of oxidant. The results are summarised in Table 2 where the
column headed R represents the ratio of CeIV consumed to the
value expected on the basis of eqns. (3) and (5), as shown
in (9). In some of these experiments the final value of [CeIV] is
R = ([CeIV]0 Ϫ [CeIV]∞)/(2[NO2Ϫ]0 ϩ [N3Ϫ]0)
J. Chem. Soc., Dalton Trans., 1999, 3311–3316
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very small, e.g. 1, 2, 7, while in others the excess of oxidant was
much larger, e.g. 9, 10, 13. The high consumption of CeIV is still
observed at lower concentrations of oxidant, comparable to
those used in the kinetics experiment. The large increase in
consumption of CeIV when azide ϩ nitrite are oxidised is clear
cut.
Writing this in the simplest form of a second order process with
(11) as the rate-determining step, and leaving aside the question
Ka
Hredϩ
red ϩ Hϩ fast
(10)
of whether (11) is better written as an inner or an outer sphere
reaction, one obtains rate law (12).
Under our conditions, with an excess of CeIV, the only stable
nitrogen compounds are N2, N2O and NO3Ϫ; these are also the
products for reactions (1) ϩ (2), (3) and (5). The two gases have
been confirmed by mass spectrometry, and nitrate has been
determined by reduction to nitrite with Cd(s), followed by
diazotisation and coupling. Some results are given in Table 3.
In these experiments we also obtained some kinetic data,
monitoring the change in absorbance A in the later stages of
reaction. The results fitted a first order rate law [Ϫ dA/dt =
k7(A Ϫ A∞)], k7/sϪ1 and the rate constants are shown in the last
column of Table 2. These k7 values are based on only a few
points, and the accuracy was reduced by problems due to evolu-
tion of gas bubbles. We cannot detect any trend in the values of
k7. This relatively slow disappearance of CeIV in the final stages
of reaction was also observed in an experiment when titrimetric
analysis was used to determine [CeIV].
k8
CeIV ϩ red
CeIII ϩ redϩ
slow
ؒ
(11)
Ϫd[CeIV]/dt = (k8/Ka)[CeIV][Hredϩ]/[Hϩ]
(12)
ϩ
The steps by which redϩ radical species such as N2H4
,
ؒ
ؒ
NH2OHϩ are converted into the final products are assumed to
be rapid and have been discussed4,6 by other workers. Values of
k8 are listed in Table 4. Values for the pKa in 0.725 mol dmϪ3
sulfuric acid are not available, and we have used literature
figures15–18 for 1 mol dmϪ3 NaClO4. This gives the two strongest
reducing agents the largest k8 values with N2H4 having a bi-
molecular rate constant not too far below the encounter limit.
This raises the question of the nature of the oxidant. In our
system calculations indicate that approximately 92% is present
as [Ce(SO4)3]2Ϫ, 8% as Ce(SO4)2, and 0.07% as [Ce(SO4)]2ϩ with
very small amounts of [Ce(OH)]3ϩ, Ce4ϩ and [Ce2(OH)2]6ϩ. Pre-
vious workers have usually sought to implicate a particular
cerium() complex as the oxidant, and it is tacitly assumed that
the more sulfate ligands bound to cerium the lower is the
reactivity, although EЊ data for the CeIV–CeIII couple are
reported19 to be almost constant from 0.5 to 4 mol dmϪ3 sul-
furic acid. The value of k8 for N2H4 is only a factor of ten below
the encounter limit, and as we calculate ca. 8% of the cerium is
present as Ce(SO4)2 this would be consistent with a diffusion
controlled reaction between Ce(SO4)2 and N2H4. Presumably,
the less sulfated complexes would also react at every encounter.
Mishra and Gupta20 also suggested Ce(SO4)2 as the active spe-
cies. However, we are sceptical that such calculations are accur-
ؒ
Discussion
The variation of k1 with [Hϩ] suggests a rate law with a 1/[Hϩ]
dependence. Morrow and Sheeres5 examined the oxidation
of hydrazine by CeIV in perchloric acid, 0.5–3 mol dmϪ3
,
and found k1 directly proportional to 1/[Hϩ]. Studies on the
cerium(IV) oxidation of hydroxylamine and of nitrous acid in
sulfuric acid4,14 showed a rate that increased as [H2SO4]total
decreased, but here the interpretation may be complicated by
the varying degrees of complexation by the sulfate ion. The
obvious explanation of the acidity dependence is that a conju-
gate base species, either of oxidant or reductant, is involved. In
the first case, this would imply [Ce(OH)]3ϩ reacting with the
reductant in its conjugate acid form red Hϩ. Morrow and
Sheeres have proposed such a mechanism involving hydrogen
transfer. It is unlikely that this could be the pattern for all four
ate enough to exclude some contribution from [Ce(SO4)3]2Ϫ
,
and would merely claim that most of the cerium() complexes
react with N2H4 at close to the encounter rate.
ϩ
ϩ
reductants. The (EЊ)Hred values for redϩ ϩ H ϩ e → Hred
ؒ
ϩ
make N2H5 and NH3OHϩ the two most powerful reductants,
yet they have the lowest reactivities. A more direct point against
the [Ce(OH)]3ϩ interpretation is that in 1 mol dmϪ3 perchloric
acid a substantial fraction of CeIV should exist as the hydroxo
complex, whereas in our system competitive complexation by
sulfate ion reduces [Ce(OH)]3ϩ to about 0.02% of total cerium
(based on Hardwick and Robertson’s formation constant3
data). However, the stoichiometric second order rate constant
for 0.94 mol dmϪ3 Hϩ in sulfuric acid is similar to the value for
1 mol dmϪ3 perchloric acid!
The same problem arises with the other reductants. Waters
and Wilson4 examined the oxidation of hydroxylamine in 3–9
mol dmϪ3 sulfuric acid and suggested an inner-sphere mechan-
ism involving Ce(SO4)2ؒNH2OH, while Treindl and Viludevá14
working in 0.5–5 mol dmϪ3 sulfuric acid suggested [Ce(SO4)ؒ
NH2OH]2ϩ. We doubt whether over such wide ranges of sul-
furic acid concentration it is possible to calculate the variation
of the relative proportions of the various complexes with suf-
ficient accuracy to identify a particular CeIV/sulfate complex as
the only active species.
Ϫ
The k8 values for NH2OH and N3 are very similar, ca. 108
The alternative interpretation is that the conjugate base
dm3 molϪ1 Ϫ1, and are significantly below the encounter limit.
s
species NO2Ϫ, N3Ϫ, NH2OH and N2H4 react with the oxidant.
Hydroxylamine is a much stronger reducing agent, (EЊ)red for
21
redϩ ϩ e → red being 0.42 V compared to 1.33 V for
ؒ
Table 3 Yields of nitrate for oxidation by 0.125 mol dmϪ3 CeIV
azide,22 but this may be offset by a much lower self-exchange
constant23 (log k11 = Ϫ12.3 compared to ϩ4.7 for azide). For
the oxidation of nitrous acid to nitric acid we have a two equiv-
[NO2Ϫ]0/mol dmϪ3
[N3Ϫ]0/mol dmϪ3
[NO3Ϫ]∞/mol dmϪ3
0.025
0.025
0.034
0.0333
0.0166
0.044
0.0166
0.0333
0.0286
0.01
0.04
0.0217
alent change. The value of EЊ for the NO2–NO2 couple22
Ϫ
1.04 V, is lower than for the CeIV–CeIII couple of 1.45 V, so
Table 4 Redox data on nitrogen substrates
Hredϩ
n
kn/dm3 molϪ1 sϪ1
pKa
k8/dm3 molϪ1 sϪ1
(EЊ)red/V
(EЊ)Hred/V
HNO2
HN3
2
3
4
5
370
2280
111
3.00a
3.5 × 105
5.9 × 107
1.2 × 108
9.3 × 108
1.04b
1.33b
0.42e
0.69g
1.23
1.61
0.77
1.16
4.44c
6.06d
7.87f
NH3OHϩ
ϩ
N2H5
13.4
a Ref. 15. b Ref. 22. c Ref. 16. d Ref. 17. e Ref. 21. f Ref. 18. g Ref. 6.
3314
J. Chem. Soc., Dalton Trans., 1999, 3311–3316

View Article Online
ϩ
the oxidation of NO2Ϫ to NO2 is reasonable. For the NO2 –NO2
could yield NH N(O)᎐NOH which could yield NH NO ϩ
᎐
2 2
NOH which would give N2 ϩ 0.5N2O ϩ 1.5H2O. We have no
evidence on this point.
couple22 EЊ is 1.51 V, and the oxidation involves a large geom-
etry change from bent to linear. Swarski et al.,24 in the course of
a study of the radiation chemistry of sodium nitrate solution in
the presence of CeIV, studied the CeIV/HNO2 reaction. They
proposed the mechanism in eqns. (13)–(16). We see no reason to
The CeIV/HNO2/HN3 system
The stoichiometric results show clearly that there is a large
excess consumption of CeIV over that expected for reactions (3)
and (5). Solutions of CeIV in sulfuric acid are very stable, and
losses by oxidation of water to O2 are negligible. This was con-
firmed by a mass-spectrometric analysis. The concentration of
nitrate ions formed is greater than the initial nitrite concen-
tration, so some of the azido nitrogen must have been oxidised.
We considered the possibility of mechanisms involving the
known intermediate N3NO formed in reaction (1) or the species
Ϫ
Ϫ
CeIV ϩ NO2
[CeIVؒNO2
]
fast
(13)
(14)
[CeIVؒNO2Ϫ] → CeIII ϩ NO2 slow
2NO2
N2O4 fast
(15)
(16)
Ϫ
N2O4 ϩ H2O
HNO2 ϩ Hϩ ϩ NO3
NNN–NO or NNN–O–N᎐O that might be formed by com-
᎐
2
bination of the radicals NO2 and N3.
modify this. The rate data of Swarski et al. when interpolated
at [H2SO4 ] = 0.725 mol dmϪ3 gave k2 = 331 dm3 molϪ1 Ϫ1, in
The rate of decomposition of nitrosyl azide, reaction (2), has
s
been found by Goldstein and Czapski28 to be >106[N3NO]
reasonable agreement with our results, assuming that their
room temperature (unspecified) was somewhat below 25 ЊC.
We have considered the possibility of an outer-sphere
mechansim. Using the simpler form of the Marcus equation25
mol dmϪ3 Ϫ1, so an encounter controlled reaction with 0.125
s
mol dmϪ3 CeIV could possibly trap N3NO before it decomposed
to N2 ϩ N2O. However, the rate of formation of N3NO is very
much less than the rates of cerium() oxidation of HNO2 and
HN3 under our conditions, so it seems unlikely that this can
account for the excess consumption of oxidant. Thus mixing
equal volumes of our stock solutions of CeIV and NaNO2 ϩ
NaN3 the calculated rates of reactions (5), (3) and (1) are 276,
48 and 0.7 mol dmϪ3 sϪ1 respectively.
22
without work terms one can use (EЊ)red and the self-
exchange constants23 for red/redϩ combined with the k8
ؒ
values for total CeIV to calculate apparent ‘self exchange con-
stants’ for CeIV/CeIII in 0.725 mol dmϪ3 sulfuric acid. Such
calculations showed no sort of self-consistency, giving values/
dm3 molϪ1 sϪ1 of 105.6 (NO2Ϫ), 109.2 (N3Ϫ) and 1015.1 (NH2OH).
These can be compared with a literature value26 of 4.2 dm3
molϪ1 sϪ1 for CeIV/CeIII in 0.4 mol dmϪ3 sulfuric acid at 0 ЊC.
We conclude that these oxidations proceed by an inner sphere
mechanism.
The molecule nitryl azide, NNN–NO2, has been made29
ϩ
Ϫ
by the reaction of NO2 with N3 in organic solvents, and is
markedly more stable than N3NO. However, it decomposes to
form dinitrogen monoxide, eqn. (18), and so although it may be
The CeIV/HNO2/N2H5؉ system
NNN–NO2 → 2N2O
(18)
The observed rate of consumption of CeIV is much faster than
the rate expected for the relatively slow CeIV/hydrazine reaction
as can be seen from Fig. 2. Comparison of the half-lifetimes
shows that the rate of consumption of CeIV is very similar to
that calculated for the rate of the hydrazine/nitrous acid reac-
tion. For [N2H5ϩ]0 = 0.035, [HNO2]0 = 0.0025, [CeIV]0 = 0.0005
mol dmϪ3 successive half-lives for CeIV are 0.0318, 00282,
0.027 s which compare with calculated values for N2H5ϩ/HNO2
of 0.027, 0.028 and 0.029 s. An exact comparison is not to be
expected because the calculated rate of the hydrazine/nitrous
acid reaction was obtained from a study10 in perchloric acid,
assuming that the rate in sulfuric acid would be the same at a
given Ho value. The results show that CeIV reacts rapidly with
an intermediate in the hydrazine/nitrous acid reaction, NH N᎐
formed by radical combination it is unlikely to be responsible
for the consumption of CeIV. The isomeric species NNN–O–
N᎐O could also be formed by radical combination, and has
᎐
been postulated as an intermediate30 to account for the prod-
ucts of the oxidation of hydrazoic acid by concentrated nitric
acid. A fragmentation reaction such as (19) would produce
NNN–O–N᎐O → N ϩ 2NO
(19)
᎐
2
nitric oxide which should be readily oxidisable to NOϩ which
would hydrate to HNO2 ϩ Hϩ. The simple form of the Marcus
equation can be used to predict a bimolecular rate constant for
an outer sphere oxidation of NO by CeIV of ca. 330 dm3 molϪ1
s
Ϫ1. Such a reaction would require 6 equivalents of CeIV to
᎐
2
NOH. There is a precursor to this species, the initial nitrosation
convert both NO molecules into nitrate and could readily
account for excess consumption of oxidant.
product NH2NHNO, and the possibility that CeIV reacts with
this before it tautomerises to NH N᎐NOH cannot be excluded.
To check whether this suggestion is reasonable we assumed
that the rate constant of combination for N3/NO2 may be
approximated as the geometric mean of the known31,32 combin-
ation rate constants of N3/N3 and NO2/NO2. A steady state
treatment based on the rates of formation and removal of N3
and NO2 leads to the prediction that the rate of combination of
N3/NO2 is ca. 1/6 of the rate of N3/N3 (which is close to the
encounter limit). Thus a reasonable assumption about the rate
of combination of N3 and NO2 predicts the formation of a
sizeable amount of the combined product, which may be, of
course, a mixture of isomers. This is, therefore, a reasonable
mechanism to account for the undoubted excess consumption
of CeIV.
The final point to discuss is the relatively slow fading of the
absorbance in the later stages of the run, which is very much
slower than would be predicted from the stopped-flow kinetics.
For example, in run 9, the half-lifetime for the consumption of
nitrite is calculated to be ca. 0.02 s, and for the reaction of azide
it should be even less. However, the absorbance changes were
measured over a period of ca. 2 min, with a half-lifetime of
᎐
2
At lower concentrations of hydrazine the rate of cerium()
consumption is somewhat less. An approximate value for the
rate constant k8 based on the assumed eqn. (17) was obtained
Ϫd[CeIV]/dt = k [CeIV][NH N᎐NOH]
(17)
᎐
9
2
by setting up a series of differential equations for d[CeIV]/dt,
d[N2H5ϩ]/dt, d[HNO ]/dt and d[NH N᎐NOH]/dt and integrat-
᎐
2
2
ing by the Gear method. A series of values for k9 was tried and
the best fit, shown in the figure, gave k9 ≈ 104 dm3 molϪ1 sϪ1
.
This fit is not perfect, and it may be that additional terms
should be included in the numerical integration. The only liter-
ature report of any related process is a study by Gupta and co-
workers27 of the cerium() reaction with hyponitrous acid
H N O . Gupta suggested oxidation of HON᎐NOH to HON᎐
᎐
᎐
2
2
2
ؒ
NO followed by a further oxidation to the acid form of
Angeli’s salt, H N O (HON(O)᎐NOH) which breaks down to
᎐
2
2
3
HNO2 and HNO, which then undergoes further reaction to
form N2, HNO3 and N2O. A similar pathway in our system
J. Chem. Soc., Dalton Trans., 1999, 3311–3316
3315

View Article Online
9 A. M. M. Doherty, K. R. Howes, G. Stedman and M. Q. Naji,
38.5 s. In run 12 the changes were observed over a 5 min time-
scale. The most likely explanation is that we are observing a rate
J. Chem. Soc., Dalton Trans., 1995, 3103.
10 J. R. Perrott and G. Stedman, J. Inorg. Nucl. Chem., 1977, 39, 325.
11 M. N. Hughes and G. Stedman, J. Chem. Soc., 1963, 2824.
12 M. N. Hughes, J. Chem. Soc. A, 1967, 902.
13 J. Fitzpatrick, T. A. Meyer, M. E. O’Neill and D. L. H. Williams,
J. Chem. Soc., Perkin Trans. 2, 1984, 927.
14 L. Treindl and A. Viludová, Collect. Czech. Chem. Commun., 1971,
36, 1654.
15 P. Lumme, P Lahermo and J. Tummavuore, Acta Chem. Scand.,
1965, 19, 617.
16 F. Maggio, V. Romano and L. Pellerito, Ann. Chim. (Rome), 1967,
57, 191.
17 P. Lumme, P. Lahermo and J. Tummavuori, Acta Chem. Scand.,
1965, 19, 2175.
18 D. Banerjea and I. P. Singh, Z. Anorg. Allg. Chem., 1967, 349, 213.
19 G. F. Smith and C. A. Getz, Ind. Eng. Chem., Anal. Ed., 1938, 10,
191.
determining fragmentation of our postulated NNN–O–N᎐O
᎐
species, followed by rapid oxidation of NO. It would be interest-
ing to see if these effects were observed for the oxidation of
mixtures of nitrite and azide by other powerful one equivalent
oxidising agents. We also considered the possibility that this
might be due to a shift in the monomer–dimer equilibrium for
CeIV as the system equilibrated after the rapid consumption of
oxidant. However, for reactions with nitrite alone as reductant,
we did not observe this effect; it occurred for reaction in a
system containing both nitrite and azide. The fact that the
slow disappearance of CeIV in the final stages of reaction was
also observed using titrimetric analysis also argues against this
interpretation.
20 S. K. Mishra and Y. K. Gupta, J. Chem. Soc. A, 1970, 2918.
21 M.-L. Hung, M. L. McKee and D. M. Stanbury, Inorg. Chem., 1994,
33, 5108.
22 D. M. Stanbury, Adv. Inorg. Chem., 1989, 33, 69.
23 D. M. Stanbury, Adv. Chem. Ser., 1997, 253, 165.
24 T. J. Swarski, R. W. Matthews and H. A. Mahlman, Adv. Chem. Ser.,
1968, 81, 164.
Acknowledgements
Thanks are due to Professor D. M. Stanbury for helpful
comments.
25 R. A. Marcus, Ann. Rev. Phys. Chem., 1964, 15, 155.
26 P. B. Sigler and B. J. Masters, J. Am. Chem. Soc., 1957, 79, 6353.
27 M. R. Goyal, P. Bhatnagar, R. K. Mittal and Y. K. Gupta, Indian
J. Chem., Sect. A, 1989, 28, 382.
28 S. Goldstein and G. Czapski, Inorg. Chem., 1996, 35, 7735.
29 M. P. Doyle, J. J. Maciejko and S. C. Busman, J. Am. Chem. Soc.,
1973, 95, 952.
30 B. M. Maya and G. Stedman, J. Chem. Soc., Dalton Trans., 1983,
2570.
31 Z. B. Alfassi and R. H. Schuler, J. Phys. Chem., 1985, 89, 3359.
32 S. E. Schwartz and W. H. White, Adv. Environ. Sci. Technol., 1983,
12, 1.
References
1 K. Clusius and E. Effenberger, Helv. Chim. Acta, 1955, 38, 1834.
2 D. L. H. Williams, in Nitrosation, Cambridge University Press,
Cambridge, 1988.
3 E. B. Robertson and H. B. Dunford, J. Am. Chem. Soc., 1964, 86,
5080; T. J. Hardwick and E. Robertson, Can. J. Chem., 1951, 29, 818,
828.
4 W. A. Waters and I. R. Wilson, J. Chem. Soc. A, 1966, 534.
5 J. I. Morrow and G. Sheeres, Inorg. Chem., 1972, 11, 2606.
6 D. M. Stanbury, Prog. Inorg. Chem., 1998, 47, 516.
7 D. W. Margerum and Z. Jia, to be published.
8 S. B. Hanna, R. R. Kessler, A. Merbach and S. Ruzicka, J. Chem.
Educ., 1976, 53, 524.
Paper 9/04080C
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J. Chem. Soc., Dalton Trans., 1999, 3311–3316