J. Chim. Phys. ( 1 998) 95, 523-535
Q EDP Sciences. Les Ulis
Influence of acidity on the oxidation rate
of nitrous acid by hydrogen peroxide
D. ~homas',J. Vanderschuren and M. Barigand
Faculte Polytechnique de Mons, Service de Genie Chimique et Biochimique,
Service de Chimie et Biochimie Appliquees,
rue de lrEpargne 56, 7000 Mons, Belgium
(Received 15 October 1997; accepted 16 December 1997)
'Correspondence and reprints
RESUME
L'oxydation de I'acide nitreux par le peroxyde d'hydrogene selon la reaction
globale:
HN02 + Hz02
HN03 + HzO
a ete etudiee dans un reacteur discontinu. Les evolutions au cours du temps des
concentrations en ions nitrites et nitrates en solution ont ete determinees par
ClectrophorQe capiflaire.
La reaction d'oxydation de HN02 par Hz02 , catalysee par les ions hydrogene
[ I ,2,3], est une reaction complexe impliquant de nombreuses etapes.
A 20°C, dans la gamrne des concentrations ([H30']:6.7 10-~-2.4
ionsA; [H202]:
4.3
.3 1 0-2 mol~l;[HN02]: 6.0 10-'-1.7 10" mow), la loi de vitesse pour le
processus d'oxydation peut s'exprimer sous la forme:
+
qui est en concordance avec les resultats des etudes precedentes.
One valeur moyenne de la constante cinetique k a ete trouvee, pour toute la gamme
de pH couverte ici, egale a 3012 12/mo12/s.
mots-clbs: oxydation, acide nitreux, peroxyde d'hydrogene, solutions acidifiees
ABSTRACT
The oxidation of nitrous acid by hydrogen peroxide according to the global reaction:
HN02 + H202 -+ HNO3 + IIzO
has been studied in a batch vessel. Variations with time of nitrite and nitrate ions
concentrations have been determined by capillary electrophoresis.
The oxidation of HNO2 by H202is known to be catalyzed by hydrogen ions [1,2,3]
and is a very complex reaction involving numerous steps.
524
D. Thomas et a/.
At 20°C, in the range of concentrations investigated in this work ([H30']: 6.7 10-'-2.4
10" ionsll; [H202]: 4.3 10-'-1.3 1v2mom; [HN02]: 6.0 10-'-1.7 lo-' molll), the rate
law for the oxidation process was found to be of the form:
which is consistent with previous studies.
An average value of the kinetic constant k has been found for all the pH range
covered here: k = 3012 12/mo12/s.
key words: oxidation, nitrous acid, hydrogen peroxide, acidified solutions
INTRODUCTION
The oxidation of nitrous acid by hydrogen peroxide according to the global reaction:
HN02 + H202
+
HNOq + H20
(1)
was found [4] to play a major role in the absorption of mixtures of NOx both in
alkaline solutions and in acidified solutions containing H202. The knowledge of the
rate of oxidation is therefore essential for the design of NOx absorption towers fed
with acidified solutions of H202.
The present work aims at securing further information on the reaction between
hydrogen peroxide and HN02 in an acid medium, at confirming the following rate law:
and at determining the orders of reaction a, h and c and the kinetic constant k.
The reaction takes place very rapidly in acid solutions, but the rate is known to
decrease markedly as the acidity is reduced. However it has been possible to examine
kinetic features in solutions of increasing acidity over a limited range of
concentrations.
Very few authors have so far investigated the kinetics of HNOz oxidation by
hydrogen peroxide and the published information is very little detailed. As a matter of
fact, this reaction is known to proceed through a sequence including peroxynitrous
acid as an intermediate species [ 1,3].
INFLUENCE OF ACIDITY ON THE OXIDATION RATE OF HN02 BY HZ02
525
Hal@enny and Robinson [I] observed that u ~ d e rneutral conditions hydrogen
peroxide and sodium nitrite do not react, but on acidification with a mineral acid, the
reaction is immediate and extremely fast, making their kinetic study more difficult.
Measurements were made in solutions containing sodium nitrite, hydrogen peroxide
and sulfuric acid. The kinetics of the oxidation reaction was estimated by decomposing
the residual hydrogen peroxide, after known time intervals, with manganese dioxide,
and by measuring the oxygen evolved. In the work of Benton and Moore [3], the
kinetic study was performed in perchloric acid solutions by the stopped-flow method.
More precisely, they investigated the formation of peroxynitrous acid (HOONO) by:
HNOz + Hz02
-+ HOONO + Hz0
(3
All available quantitative results are reported in Table I.
TABLE I: Results of previous studies relative to oxidation rate of HNOz by H z 0 2
EXPERIMENTAL
Experiments were achieved in a 12 cm ID perfectly stirred glass batch vessel with a
total volume of approximately I litre. Temperature, which is an essential parameter, is
regulated by a thermostatic bath and controlled by thennometers.The H202
concentration is determined bv means of an iodometric method while the
D. Thomas et a/.
526
concentrations of nitrite and nitrate ions are measured by electrophoretic capillary ion
analysis (CIA).
Because of the rate of the reaction, small quantities of reagents were used, involving
small quantities of products. Large dilutions are therefore necessary to reach the low
concentrations required. An error of analysis lower than 1% can be expected with the
CIA in the range of concentrations investigated in this work.
Nitrous acid is a weak acid. The value of the dissociation constant K,, used in our
work is 5.2 lo4 [5] at all ionic strengths, corresponding to pKd = 3.28 at 20°C.
For each set of kinetic experiments, fiesh solutions of HN02 were prepared by
acidifying a solution of sodium nitrite with HCI. The oxidation reaction is initiated by
introducing a given quantity of the H202solution in the reactor containing the mixture
of NaN02 and HCI. At the different times the reaction medium had to be analysed, a
small given quantity of the solution was taken and then immediately injected in a given
quantity of NaOH. This operation aimed at making the sample alkaline and stopping
the oxidation reaction, as H202does not oxidize nitrite ions [I]. Moreover in these
alkaline conditions, the ion capillary analysis is applied to the determination of the
total nitrite concentration which represents the sum of HN02 and NOz-concentrations
initially contained in the acid sample:
mo2-ltOt
= [HNO~I
+1 ~ 0 ~ 1
(4)
Consequently a model is necessary to compute [HNOz], [NOz'] and [H,O'] at
equilibrium in the acid reaction medium, from [NOilt, measured in alkaline medium.
These three concentrations required to characterize the complete composition of the
solution are calculated by solving the set of three equations:
- mass balance of nitrous compounds:
BaNO2] = [HN02]+ WO2-1 + fNo<] = lY02-1t0t + WO<l
- electroneutrality relation:
ma"]
+ [H~O"]= [CI-] +- [NOT] + [NOi]
- dissociation equilibrium of nitrous acid:
J. Chirn. Phys.
INFLUENCE OF ACIDITY ON THE OXIDATION RATE OF HNO;! BY H202
Kd =
[H30f].[N02-]l[HN02]
for reaction: H3O' + NOi o HNO:! + Hz0
from known values of I& and of the concentrations of HCI and
527
(7)
(8)
m027t,l whch,
at
initial time, corresponds to maN02 J. This model includes the following assumptions:
[Na+]=[NaNOz],[CI'] =[HCl] and [OH-] negligible.
The following figue (Fig.1) illustrates some of the results yielded by our model:
calculated values of the equilibrium concentrations of HN02, NOY and of the pH of
the solution are given versus cologarithm of the initial concentration of HCI, and for
different initial concentrations of NaNO? dissolved in the solution.
Figure I: Variation, with the irritial concentration of HCI. ofthe calculated eqt~ilibrii~m
concentratii>tlsand of the pH, for different initial concentrationsQf NahQ
It can be seen that a substantial part of the hydrogen ions are consumed by nitrite
ions, in order to form HN02, as NaN02 is introduced in excess relatively to HCI.
Therefore the calculated value of pH is higher than the cologanthrn of the initial
D. Thomas et a/.
528
concentration of HC1, i.e. the pH value which should have been obtained without
NaN02. The higher the excess quantity of NaN02, the more pronounced this effect.
A pHmeter placed in the system allows to compare experimental and calculated pH
values of the solution.
Due to the consumption of HNO2 for oxidation by Hz02 (reaction (1)) which is
perpetually counterbalanced by the production of HN02 thanks to the reaction at
equilibrium between N02- and H30' (reaction (8)), the oxidation rate to be measured is
equal to:
A second-order polynomial was used to describe the time-dependence of the [NOz-],<,,
decrease:
[N02-]ror = a t2 + bt +C
( 1 0)
The initial concentration method was chosen for the determillation of kinetic
parameters appearing in relation (2). Initial rates are simply equal to parameter h in
this relation as:
RESlJLTS
Oxidation rate of HN02 at pH close to 4
For the determination of orders of reaction at this pH value, we adopted the
procedure developed by Halfpenny and Robinson [I], and introduced in the reaction
medium a concentration of [NaN02] much higher than that of HCl, which involves, as
already mentioned, pH>-log[HCl].
In order to apply the initial concentration method, we needed to find a concentration
range in which one initial concentration, either [HN02Ioor [H30'Io, does not vary with
various introduced quantities of NaN02. With [HCl]=1.2 10" molA and increasing
INFLUENCE OF ACIDITY ON THE OXIDATION RATE OF HNO2 BY Hz02
529
vaNO21, we observe in Table I1 that [KNO2Iocan be considered as constant and
almost equal to [HCI].
TABLE 11 : Calculated equilibrium H30' and HNO2eoncentr~tions,
at a constant HCI
molarity, for various NaNOz concentrations
IHCII
(1O ~ O I A )
1.2
12
1.2
1.2
Relative max.
variation
[NaN021
( 1 0.' mol/l)
5
6
7
8
[H~O'IO
( I oJ ~ O M )
1.303
1.067
0.902
0.787
PH~
3.88
3.97
4.04
4.11
IT
0210
(10 rnol/l)
I0 697
10.933
11.098
39.6%
11.219
4.67%
In these conditions, the pH value is not constant but remains close to 4.
As the oxidation reaction proceeds, a typical variation with time of concentrations
of the different species in solution is shown in Figure 2, for [H202]0=8.7 lo-",
[HC1]=1.2 lo-% and waNO2]=6.9 10" M. On this graph, pOT]lol is measured by
CIA, [NOz-], Po?-1,
[HN02]and [H30'] are calculated thanks to the model developed
hereabove.
t (min)
Figure 2: Variation with time o f N & , , NO;, HN02, NO.?a d HNO? concentrations at pH clo.se
to 4 ((HzOrJo 8 . 7 I@' M: ( H ( - I / --1.2 10 M ; [NaNO2]-6.9 10.' M)
J Chim Phys.
D. Thomas et a1
530
The conditions reported in Table 11 are favourable to the determination of a. order
of reaction relative to the hydrogen ions, thanks to a variable [H30+lo with constant
[H20210and [HN0210. Therefore, equation (2) becomes for initial rates:
vo = k l . [ ~ , O t ] ; where kl= ~ . [ H ~ O .[HNO;!];
~]:
(12)
Results are given in Table 111.
TABLE 111: Determination of order a of reaction with respect to RJO' for pH close to 4
The different values of In vo when plotted against corresponding values of In f i 0 + / 0
exhibit a straight line with a slope of 0.93, indicating an initial rate directly
proportional to the hydrogen ions concentration.
The effect of hydrogen peroxide concentration on the reaction rate was studied by
measuring conversions of HN02 in NO?, after various times, in solutions containing
fixed initial concentrations of NaNOz (6.9 10" M) and of HCI (1.2 10'~M), then also
of HNOz ,but various initial H202concentrations (4.43 lo-', 8.82 1o", I. l l 10-2 and
1.3 1
M). In these conditions, expression (2) reduces to:
b
~o=kz.[H~O~]~
(13)
Figure 3 exhibits the linear rclationship obtained when plotting In vo versus In @202/0.
The slope is equal to 0.97 and an order b of reaction is observed with respect to
hydrogen peroxide equal to one.
With a given excess of NaN02 (=5
HzOz(-8.82
mom), for a given initi'al concentration of
M) but increasing initial concentrations of HCI (9.17 104, 1.23 lo-',
1.54 lo-' and 1.83 10" M) , two quantities increase simultaneously, as can be seen on
Figure 1, i.e. [HNO2I0and [H3OiIo,and relation (21 becomes:
INFLUENCE OF ACIDITY ON THE OXIDATION RATE OF HN02 BY H202
vo= k~.(l#z0+lo.[~-M0210)c
531
(14)
Figure 4 shows, on a log-log plot, the results of the experiments in terms of In vo as a
h c t i o n of In (fl-?O'Jo./HN02/0).A straight line was obtained, with a slope of 0.99.
As a, determined hereabove, is equal to unity, the value of c was found to be 1 too.
Figure 3: Delermination of order b of
Figure 4: Ueterminationof order c of
reaction with respect lo H N 0 2 (pH close to -I)
reffclionwifh reSPecf to H 2 0 2 (pH clo-~eto ?')
Oxidation rate ofHNOz at pH equal to 3
In our study of the oxidation rate at a pH of 3, given the increase of the hydrogen
ions concentration, we had to work with lower concentrations of NaN02, of about 10"
moV1, in order to keep measurable reaction rates.
Variations with time (Figure 5) of concentrations in a solution are totally different from
those obtained at pH close to 4 (Figure 2): the pH value can be, in this case,
considered as constant during all the oxidation reaction.
The hydrogen ions concentration and the kinetic constant k can be put together in the
relation (2) which becomes then:
v = k ? [ ~ ~[ O
H A~ Q] ~] ~with
k' = k.[H30+Ia
(15)
For the determination of kinetic parameters, we used the initial concentration
method, as concentrations of HNOz and H202 vary with time. By varying initial
concentrations of Hz02 (3.5 10'-1.3 lo-' moM), at constant initial concentrations of
HN02, and by varying initial concentrations of HNO2 (6.8 10-~-1.4loJ moVI), at
J. Chim. Phys.
532
D. Thomas et a/.
constant initial concentrations of HzOz, it is possible to determine the orders of
reaction relative to H202
(h) and HNOz ( c )respectively.
t (min)
Figure 5: Variation with lime o j / i V 0 2 J,,,,, /HN02] and WOiJ at pH rquul to 3
( [ N c N O J 9.5 I O S M; (Hz()& 3.44 1 0 .M;
~ p/(:lJ 1.22 1 0 3M)
the following expressions are derived fiom (1 5):
-
- at constant [HN02]o: vo k', [H ,02I$
(16)
- at constant [H202]0: vo= k'Z [HNOZ]
(17 )
Globally, fiom the 16 experiments we carried out (Table IV), with the same procedure
than at pH close to 4, we have concluded orders of reaction h and c equal to 1.
TABLE IV: Initial rates and determination of orders b and c of reaction for pH=3
(concentrations rre in mol~land reaction rates in mows)
INFLUENCE OF ACIDITY ON THE OXIDATION RATE OF HN02 BY H202
533
Kinetic determinations at pH=2 andpH=l.6
Experiments were first carried out at these two pH values with similar procedure, to
determine orders of reaction, than at pH 4 and 3. In our attempt to investigate the
~nfluenceof increasing acidities on the kinetics of oxidation, test runs were possible up
to a pH value of 1.6 corresponding to an operating limit with too important
interferences with CI- ions, in the capillary ion analysis.
For all our results, we concluded that at both pH values, the reaction rate constants
b and c were equal to 1.
Order of reaction with respect to HJO+and kinetic constant
We are now able to determine the order a of reaction relative to hydrogen ion, on
the whole pH range (1.6-4.1) covered in our work.
Actually, for all the pH investigated, an average value of k' (see d e b t i o n in equation
( 1 5)) could be determined.
Figure 6 gives a representation of the law:
and leads to a=0.94, which can be assimilated to 1
Figure 6: Determination of order a of
reaction with respect 10 the hydroget1 ion
-
fiptre 7: ~eterminbtionqf the kinetic
constant of the reaction
The following step consisted to plot k' values versus the concentration of hydrogen
ions (see Figure 7). The slope of the straight line obtained corresponds to the average
D.Thomas eta/.
5 34
value of the kinetic constant k: k=3012 12/mo12/s.This constant varies within the error
associated with the method of analysis (dilutions, CIA) and the method of
determination (least-squares technique): +8%.
Table I allows to compare this value with those obtained by different authors. This
comparison gives the same orders of magnitude.
CONCLUSIONS
The kinetics of oxidation of nitrous acid by hydrogen peroxide according to the
global reaction HNO2 + Hz02
+ HNOT + H20
was studied at 20°C in a small
laboratory batch vessel, in presence of different concentrations of HCl.
The initial concentration method was applied for this study in which the
variation with time of nitrate ions in solution was determined by the capillary ion
analysis.
In the range of concentrations investigated in this work ( [H30']: 6.7 10"-2.4
1
ions/l; [H202]: 4.3 10.~-1.310.' moM;
[HN02]:
6.0 10-~-1.710.' moVl), the rate
law for the oxidation process was found to be:
with k=30 12 12/mo12/s.
These results are consistent with those reported in previous studies.
ACKNOWLEDGEMENTS
The authors acknowledge for the support of the National Fund for Scientific
Research of Belgium which gave a fellowship to DT and the h d s for the purchase of
the electrophoretic ion capillary analyser .
J. Chim. Phys.
INFLUENCE OF ACIDITY ON THE OXIDATION RATE OF HN02 BY H202
REFERENCES
./otlmal~x1 Halfpenny E, Robinson PL (1952) .J Chem Soc (A) ,928-938
2 Anbar M, Taube H (1954) J Am ('hem Soc 76,6243-6247
3 Benton DJ, Moore P (1970) J Chem Soc (A), 3 179-3 182
4 Thomas D, Vanderschuren J (1997) Ind Eng Chem Res 36(8), 33 15-3322
5 Vassian EG, Eberhardt VH (1958) .J Phys Chem 62,84-87
J. Chim. Phys.
535