Critical phenomena in acetone oxidation by nitric acid

Article abstract of DOI:10.1007/s11172-008-0280-8

The kinetic regularities of acetone oxidation by aqueous nitric acid solutions (5.86-58.31 wt.%) were studied using a differential automatic microcalorimeter. The critical phenomena were discovered, which manifest themselves as a abrupt change in the initial heat release rate at a minor change in the temperature or acid concentration. The abrupt change in the oxidative activity of the reactant at a minor change in the system parameter was assumed to be related to changes in the structure of the solution and, as a consequence, in the solvation energy of the reactants at a certain acid/water ratio in the solution.

Full text of DOI:10.1007/s11172-008-0280-8

Russian Chemical Bulletin, International Edition, Vol. 57, No. 10, pp. 2065—2071, October, 2008
2065
Critical phenomena in acetone oxidation by nitric acid
Yu. I. Rubtsov,† A. I. Kazakov, T. V. Sorokina, and G. B. Manelis
Institute of Problems of Chemical Physics, Russian Academy of Sciences,
1 prosp. Akad. Semenova, 142432 Chernogolovka, Moscow Region, Russian Federation.
Fax: +7 (496) 524 9676. Eꢀmail: akazakov@icp.ac.ru
The kinetic regularities of acetone oxidation by aqueous nitric acid solutions (5.86—58.31 wt.%)
were studied using a differential automatic microcalorimeter. The critical phenomena were
discovered, which manifest themselves as a abrupt change in the initial heat release rate at a
minor change in the temperature or acid concentration. The abrupt change in the oxidative
activity of the reactant at a minor change in the system parameter was assumed to be related to
changes in the structure of the solution and, as a consequence, in the solvation energy of the
reactants at a certain acid/water ratio in the solution.
Key words: kinetics, mechanism, oxidation, acetone, nitric acid, critical phenomena.
The reactions of nitric acid with organic compounds,
mainly nitration reactions, which are widely used for synꢀ
thesis of dyes, drugs, explosives, gunpowders, etc., are
being intensively studied in the modern organic chemisꢀ
try. Nitration is almost always accompanied by a parallel
oxidation reaction, whose rate under normal conditions
is much lower, but the thermal effect is substantially
higher. The change in the parameters of the process may
result in an abrupt increase in the heat release rate in the
system and in the loss of the thermal stability of the proꢀ
cess. According to the current knowledge, the kinetics and
mechanism of liquidꢀphase oxidation of different classes
of organic compounds including ketones with molecular
oxygen can be described by the regularities of chain reacꢀ
tions with degenerate chain branching.1 In turn, data on
the kinetics and mechanism of ketone oxidation by bound
oxygen, particularly, by the systems based on nitric acid
solutions, are rather scarce, and the available works are
mainly preparative. It was found2—6 in a series of studies
on the liquidꢀphase oxidation of aliphatic hydrocarbons by
nitric acid solutions that the process under study is classiꢀ
fied as a reaction where acidꢀbase equilibria play a subꢀ
stantial and in many aspects determining role along with
the radical steps characteristic of oxidation. A combination
of radical irreversible and ionꢀmolecular reversible steps is the
main specific feature of oxidation by systems based on nitric
acid solutions. This opens new possibilities to control the rate of
the overall process by shifting the acidꢀbase equilibria.
c 34.7 mol.%), and their concentrations were determined by
acidometric volumetric titration. The heat release rate during
oxidation was measured with a differential automatic calorimeter7
in glass sealed ampules without cold parts, which made it possible
to retain all reaction components in the reaction bulk. The
temperature interval studied was 17—99.4 °С, and the concentraꢀ
tion of HNO3 solutions expressed as a molar ratio of water to
acid (nH /nHNO ) was 2.50—56.14. The investigated acetone—
2O
aqueous nitric ac3id system is homogeneous in the region of the
studied concentrations of nitric acid solutions. A high HNO3
excess with respect to acetone was used in order to neglect the
change in its content during the process.
Results and Discussion
We studied the nature of critical phenomena in the
acetone—aqueous nitric acid system. It was found by the
integration of the time dependence of the experimental
heat release rates that the thermal effect of acetone oxidaꢀ
tion (Q0) is ∼770 kJ mol–1, which corresponds to the oxiꢀ
dation mainly to carbon dioxide with a minor admixture
of acetic acid. Acetone is oxidized to the same products
during liquidꢀphase oxidation by molecular oxygen.1 The
fraction of nitration products among the reaction prodꢀ
ucts should be insignificant, because the estimated heat of
acetone nitration is ∼120 kJ mol–1 of acetone.
CH3C(O)CH3 + HNO3 = CH3C(O)CH2NO2 + H2O
If we assume that the decrease in the experiꢀ
mental heat compared to the theoretically possible value
(800 kJ mol–1) is related to the contribution of nitration,
then at most 5 mol.% acetone should be consumed by this
reaction to yield corresponding nitration products.
The experimental temperatures and concentrations
of HNO3 solutions expressed in the number of moles of
Experimental
Acetone (special purity grade) was used in experiments.
Solutions of nitric acid were prepared from HNO3 (reagent grade,
† Deceased.
Published in Russian in Izvestiya Akademii Nauk. Seriya Khimicheskaya, No. 10, pp. 2028—2034, October, 2008.
1066ꢀ5285/08/5710ꢀ2065 © 2008 Springer Science+Business Media, Inc.

2066
Russ.Chem.Bull., Int.Ed., Vol. 57, No. 10, October, 2008
Rubtsov et al.
1
2
T/°C
dQ
dt
___
/ W g–1
100
80
2.5
1
2.0
1.5
1.0
II
60
40
I
III
20
0
10
20
30
40
50
nH O/nHNO
2 3
2
Fig. 1. Temperatures and concentrations of HNO3 at which
acetone oxidation was studied: I, selfꢀaccelerating reaction, the
presence of NO2 in the products (1); II, selfꢀaccelerating reacꢀ
tion, NO2 is not visually observed in the products (1); III, the
heat release rate is lower than the microcalorimeter sensitivity
(∼0.01 mW) during the observation time up to several days (2).
3
4
5
6
7
8
water per mole of the acid at which the oxidation experiꢀ
ments were carried out are shown in Fig. 1. The whole
interval of experimental temperatures and concentrations
of nitric acid solutions can be divided into three regions
depending on the observed regularities. In region I of
relatively high concentrations of HNO3 solutions with
0
400
800
1200
t/min
Fig. 2. Plots of the heat release rate vs time during acetone
oxidation by HNO3 solutions of different concentrations at 73.8 °C,
nH /nHNO : 20.98 (1), 24.61 (2), 27.06 (3), 30.31 (4), 35.69 (5),
2O
40.69 (6), 436.72 (7), and 56.23 (8).
nH /nHNO < 4 and low temperatures (Т < 40 °С), the
2O
3
reaction occurs with selfꢀacceleration, the curve of the
heat release rate is characterized by an induction period
followed by the evolution of the maximum amount of
heat in a narrow time interval. The reaction rate depends
monotonically on the temperature and concentration of
HNO3 solutions, and the accumulation of NO2 is visually
observed during the reaction.
rimeter during several days. The corresponding region in
Fig 1 is designated as region III. For example, at Т = 34.7 °С
and the nitric acid concentration nH /nHNO = 3.97 (reꢀ
2O
3
gion I) the maximum oxidation rate is achieved for ~250 min,
whereas at nH /nHNO = 4.31 (region III) no heat release
2O
3
is observed during 38 days (Fig. 3). Upon further decrease
in the initial concentration of the nitric acid solution (at a
constant temperature) at the next increment correspondꢀ
ing to the transition from region III to region II, the
oxidation process can be detected again with the times of
achievement of the maximum rate about several hours.
Such critical phenomena as two abrupt changes in the time
of achievement of the maximum rate (from hours to days
and then again from days to hours) with the consecutive
decrease in the initial concentration of the nitric acid
solution were observed at 47.2, 51.6, and 55.7 °С. The heat
release curves obtained at 55.7 °C are shown in Fig. 4.
In region III, as in region II, no NO2 is visually obꢀ
served among the products. The boundary between reꢀ
gions I and III, determining the critical concentration
and temperature values, can be approximated by a linear
equation (Fig. 1)
In region II of relatively low concentrations of HNO3
solutions (nH /nHNO > 10) and relatively high temperaꢀ
2O
3
tures (Т > 60 °С), the oxidation also occurs with selfꢀ
acceleration, the maximum rate is achieved within several
hours, but the release of the maximum amount of heat
occurs in a larger time interval than that for the process in
region I. The heat release curves of the oxidation process
in region II are shown, as an example, in Fig. 2. As in
region I, the reaction rate varies regularly with tempeꢀ
rature and concentration of HNO3 solutions, and no
accumulation of NO2 in the reaction products is visually
observed.
When the concentration of an HNO3 solution deꢀ
creases with an increment of 1—2 wt.% in region I at a
constant temperature, the reaction rate decreases smoothly,
but at some next insignificant change in the concentraꢀ
tion of an HNO3 solution the process rate decreases
jumpwise to the value that is not detected by a microcaloꢀ
Тcr = 13.3•(nH O/nHNO )cr – 18.8.
3
2

Critical рhenomena in acetone оxidation
Russ.Chem.Bull., Int.Ed., Vol. 57, No. 10, October, 2008
2067
dQ
dt
dQ
dt
___
___
/ W g–1
/ W g–1
9
1
4
1
6
3
2
1
2
2
10
3
4
2
4
3
5
4
0
400
6000 t/min
5
6
7
Fig. 3. Plots of the heat release rate vs time during acetone
8
0
oxidation by HNO3 solutions of different concentrations at
100
500
1000 1500 t/min
34.7 °C, nH O/nHNO : 2.88 (1), 3.11 (2), 3.38 (3), 3.97 (4), and
2
4.31 (5); the heat rel3ease rate is lower than the microcalorimeter
Fig. 4. Plots of the heat release rate vs time during acetone
oxidation by HNO3 solutions of different concentrations at
51.7 °C, nH O/nHNO : 6.05 (1), 7.09 (2), 10.48 (3), 13.22 (4),
sensitivity.
2
3
14.68 (5), 16.73 (6), 18.71 (7), 20.98 (8), 6.74 (9), and 6.74 (10).
The processing of the experimental curves by the
Origin program showed that the oxidation rate in regions I
and II is described by the empirical equation
concentration of undissociated and unbound nitric acid
in aqueous solutions8 are presented in Figs 5 and 6.
The abrupt jumpwise change in the apparent rate
constant in the initial step at a consecutive transition
from region I to region III and then from region III to
region II (see Figs 5 and 6) corresponds to the critical
m
dη/dt = k1(1 – η)n + k2(1 – η)n•η ,
where
is the conversion of the reacꢀ
tion, dη/dt = 1/Q0•(dQ/dt) is the process rate, dQ/dt is
phenomena. In region I at different nH /nHNO ratios the
2O
3
apparent activation energy for the initial step is approxiꢀ
mately the same (∼200 kJ mol–1) and the preꢀexponential
factors increase with increasing acid concentration as
follows:
the heat release rate, and
tion heat.
is the total oxidaꢀ
The values of the k2 rate constant and the reaction
orders were determined by the processing. It was found
that in region I nav = 2 and mav = 1, and in region II the
corresponding values are 2 and 0.5. The k1 rate constants
are too low to be determined with the Origin program.
They were found by the comparison of the experimental
kinetic curves η(t) with the curves calculated using speciꢀ
fied values of k1, k2, n, and m by a numerical calculation
procedure developed at the Institute of Problems of
Chemical Physics (Russian Academy of Sciences). The
program is based on an implicit firstꢀorder difference
scheme with Newton iterations. The dependences of the
reaction rate constants k1 and k2 on the temperature and
nH O/nHNO
Preꢀexponential factor (min–1
)
2
3
56.22
20.98
10.48
6.6•1022
3.1•1027
7.2•1029
The apparent rate constants k2 in the region of inꢀ
creasing rates are characterized by smooth (no jumps)
dependences on the temperature and nitric acid concenꢀ
tration in the entire interval of external parameters. The
apparent activation energy for this stage is approximately
the same for the oxidation by solutions of different conꢀ

2068
Russ.Chem.Bull., Int.Ed., Vol. 57, No. 10, October, 2008
Rubtsov et al.
lg k1/min–1
lg k2/min–1
1.5
a
b
0
–5
6
5
1.0
0.5
1
2
3
4
5
0.0
4
–10
–15
–20
–25
–0.5
–1.0
–1.5
–2.0
–2.5
3
1
2
1
T
__
/K–1
0.0028
0.0030
0.0032
1
T
0.0028
0.0030
0.0032
__
/K–1
Fig. 5. Plots of logk1 (a) and logk2 (b) vs 1/Т during acetone oxidation by nitric acid solutions of different concentrations,
nH /nHNO : 56.22 (1), 20.98 (2), 10.48 (3), 5.54 (4), and 3.68 (5).
2O
3
lg k1/min–1
lg k2/min–1
a
b
0
–1.0
–10
1
–20
–30
–40
–50
2
3
4
–1.5
–1.0 lg CHNO /mol L–1
–2.0
–1.5
–1.0 lg CHNO /mol L–1
–2.0
–1.5
3
3
Fig. 6. Plots of logk1 (a) and logk2 (b) vs logCHNO during acetone oxidation by nitric acid solutions at different temperatures, T/°С:
3
55.7 (1), 51.6 (2), 47.2 (3), and 34.7 (4).

Critical рhenomena in acetone оxidation
Russ.Chem.Bull., Int.Ed., Vol. 57, No. 10, October, 2008
2069
centrations (∼120 kJ mol–1), while the preꢀexponential
factors change as follows:
The critical phenomena are an important feature of
the acetone oxidation by nitric acid solutions. They
appear as a sharp change in the heat release rate at an
insignificant change in the temperature or acid conꢀ
centration. The critical phenomena observed in chain
branched reactions at an insignificant change in the
parameter of the system are related to the abrupt change
in the apparent rate constant in the selfꢀacceleration step
due to the difference in the rate constants and orders of
the chain branching and termination reactions.9 In the
case of acetone oxidation in solutions by bound oxygen
of nitric acid, the critical phenomena, probably, have
another nature. The smooth change in the parameters of
the system (temperature or concentration) is accompaꢀ
nied by a jumpwise change in the initial rate of the proꢀ
cess, which can be associated to the same kind of changes
in the concentration of the oxidized compound or the
oxidizing reagent in the initial step. For acetone oxidaꢀ
tion by nitric acid solutions, the apparent rate constant
for the selfꢀacceleration step exhibits a smooth depenꢀ
dence on the temperature and HNO3 concentration in
the entire interval of concentrations of HNO3 solutions,
including the initial parameters at which critical phenomꢀ
ena are observed (see Figs 3 and 5). This is the basic
difference between the nature of critical phenomena
observed previously in the branched chain reactions and
in the present study.
nH O/nHNO
Preꢀexponential factor (min–1
)
2
3
56.22
20.98
10.48
5.54
3.68
2.92
1.7•1014
1.3•1015
1.9•1017
9.0•1017
2.9•1019
2.2•1020
As follows from the slope ratios of the straight lines,
the apparent rate constant for the initial stage k1 is proꢀ
portional to the concentration C of undissociated and
unbound nitric acid, and the apparent rate constant k2 in
the selfꢀacceleration stage is proportional to C0.75 (see
Fig. 6).
The influence of sodium nitrite additives on the rate of
acetone oxidation was studied at a fixed concentration of
HNO (nH /nHNO = 10.17, region III) and 33.5 °С. The
2O
3
3
added amount of nitrite ranged from 0.05 to 0.9 mole per
mole of acetone. The sodium nitrite additives substanꢀ
tially increase the initial rate. The initial rate was found to
be proportional to the initial concentration of sodium
nitrite С0 – to the 0.5 power
NO2
–
NO
NO
0.5
С0
–•106
(dQ/dt)t=0•10–1/(С0
)
2
2
/mol L–1
/W•L0.5•g–1•mol–0.5
Throughout studied interval of solution concentraꢀ
tions, except for region II, the initial rate of acetone oxiꢀ
dation is proportional to the concentration of undissociꢀ
ated and unbound acid. It can be assumed that, as for the
earlier studied reaction of hydrocarbon oxidation by
nitric acid,1 the reaction rate is controlled by the rate
constant and equilibrium concentration of molecular
HNO3. The concentrations of other potential oxidants
present in the equilibrium system (nitric acid anhydride
N2O5 and the product of its ionization NO2+) are substanꢀ
tially lower than the concentration of molecular HNO3.
Initially, there is no nitrogen dioxide in the system,
because the relatively dilute solutions of nitric acid used
are stable at the experimental temperatures.10 It can be
assumed that the observed jumpwise changes in the initial
rate of acetone oxidation on going to region II are assoꢀ
ciated with an abrupt change in the solvation energy of
acetone or oxidizing agent at a certain water/acid ratio
in solution at a given temperature. There exist dynamic
equilibria between diffusionꢀaveraged supramolecular
structures of hydrated proton and anion in nitric acid
solutions, with singularities at particular temperatures
and concentrations.11 The specific features of the superꢀ
molecular structure and specificity of incorporation of an
acetone molecule into these structures in a certain interꢀ
val of concentrations and temperatures can result in a
lesser accessibility of the acetone molecule to the attackꢀ
ing nitric acid molecule and, correspondingly, in an abrupt
111
100
19
17
8.1
6.5
6.3
6.5
6.4
6.5
The k1 and k2 values were determined at a fixed soꢀ
dium nitrite additive in an amount of 1.15•10–4 mol L–1
in HNO with nH /nHNO = 10.17 in the temperature
2O
3
interval from 33.5 to 54 °С3
Т/°С
k1•103/min–1
k2•103/min–1
33.5
38.3
39.3
43.3
44.5
48.8
54.0
0.2
1.3
1.6
1.9
1.4
3.6
7.8
0.5
1.3
1.7
2.9
2.3
5.0
10.8
The temperature dependence of k1 can be presented as
4.5•1017exp(–124•103/RT) min–1. The activation enerꢀ
gies in the equations for k1 and k2 are equal to each other
within the limits of experimental error, while the rate
constants k2 are on the average 1.5 times higher than k1.
The oxidation with a sodium nitrite additive proceeds
with a substantially lower rate constant k2 compared to
the systems without additives.

2070
Russ.Chem.Bull., Int.Ed., Vol. 57, No. 10, October, 2008
Rubtsov et al.
•
CH3C(O) + NO2 → CH3C(O)ONO,
(8)
(9)
decrease in the initial oxidation rate. Interrelations
between the structure of nitric acid solutions and kinetic
phenomena will be a subject of future studies.
CH3C(O)ONO + H2O → CH3COOH + HNO2,
The initial step of acetone oxidation can be either the
abstraction of an H atom from the methyl group of
the ketone form of acetone by the HNO3 molecule
•
CH2O + NO2 → CHO + HNO2.
(10)
The oxidation by HNO3 solutions differs from the
oxidation by molecular oxygen, in that the concentration
of NO2 leading the process can increase following the
acidꢀbase equilibrium reaction (1). For two NO2 molꢀ
ecules consumed in reactions (5) and (6), three NO2
molecules are regenerated via equilibrium (1). The NO2
molecules formed can again enter into the oxidation reꢀ
action to form NO and HNO2. This process can formally
be considered as a branched chain reaction. Thus, when
nitrogen oxide is accumulated in the system, reaction (1)
can be one of the main sources of the fast increase in
the NO2 content and of strong selfꢀacceleration of the
oxidation of organic compounds by nitric acid. According
to equilibrium (1), the NO2 concentration in the system
should be proportional to the concentration of free nitric
acid molecules to the 0.66 power, which is close to the
experimentally observed order of the apparent rate conꢀ
stant k2 with respect to the concentration of free nitric
acid molecules. The role of chain termination reactions is
played by the oxidation by NO2 with the formation of N2
•
CH3C(O)CH3 + HNO3 → CH3C(O)CH2 + H2O + NO2,
or the addition of the HNO3 molecule to the double bond
in the enolic form of acetone followed by the reduction of
nitric acid to nitrous acid and acetone oxidation to the
αꢀoxy compound.12 The oxidation rate by HNO3 molecules
is low and controls only the initial rate of oxidation.
Nitrogenꢀcontaining molecules with the oxidation
state of lower than +5 (NO2, N2O4, N2O3, HNO2) formed
upon the reduction of HNO3 are stronger oxidants than
molecular HNO3. In HNO3 solutions, their concentraꢀ
tions are determined by mobile equilibria. On going from
concentrated HNO solutions (nH /nHNO = 2.3) to diꢀ
2O
3
3
lute solutions with increasing the water content the equiꢀ
librium shifts from NO2 to N2O3, then to HNO2 and,
finally, to NO2– according to the equations13
3 NO2 + H2O
NO + NO2
NO + 2 HNO3,
(1)
(2)
(3)
(4)
N2O3,
•
CHO + NO2 → CH(O)ONO,
(11)
(12)
(13)
N2O3 + H2O
HNO2 + H2O
2 HNO2,
CH(O)ONO → CO2 + HNO,
H3O+ + NO2
.
–
2 HNO + HNO2 → N2 + H2O + HNO3.
As the process develops and these reduced nitrogenꢀ
containing molecules appear in the system, the oxidation
rate in the developed process is controlled by the reaction
with that oxidant whose oxidative activity, determined by
the oxidation rate constant and the equilibrium concenꢀ
tration, is the highest in the studied interval of HNO3
concentrations. In region I, where the accumulation
of NO2 is visually observed, the rate of the developed
acetone oxidation of acetone is determined by its reaction
with NO2 and increases with increasing the equilibrium
concentration of the latter.
When equilibria (1)—(4) shift to the left, NO2 is
formed, and these reactions result in chain branching. The
chain propagation reactions,14 in which NO2 is consumed
to form the product of acetone oxidation (MeCOOH), are
presented below
During acetone oxidation by HNO3 solutions, the recomꢀ
bination of NO2 with CH3C(O)C•H2 and CH3C(O)CH2O•
radicals can probably lead to the formation of the nitraꢀ
tion product. The reactions of formation of the nitro deꢀ
rivatives and nitric acid esters and the equilibria (1)—(4)
along with the equilibrium
–
2NO2
N2O4
NO+ + NO3
shifted to HNO2, N2O3, N2O4, NO+, and NO2 are also
chain termination reactions when the oxidation is led by
nitrogen dioxide.
–
As in region I, in region II (HNO3 solutions with
nH /nHNO > 10 and relatively high temperatures), the
2O
3
reaction rate shows a smooth dependence on the temꢀ
perature, concentration of the HNO3 solutions, and
conversion with no critical jumps. It was established for
this region of HNO3 concentrations that the initial oxidaꢀ
tion rate is proportional to the concentration of sodium
•
CH3C(O)CH3 + NO2 → CH3C(O)CH2 + HNO2,
(5)
(6)
(7)
•
•
CH3C(O)CH2 + NO2 → CH3C(O)CH2O + NO,
nitrite added (С0 –). This means that NO2 is the leading
oxidizing agent in region II both in the initial step in the
presence of the nitrite ion and for the developed process
NO
2
•
•
CH3C(O)CH2O → CH3C(O) + CH2O,

Critical рhenomena in acetone оxidation
Russ.Chem.Bull., Int.Ed., Vol. 57, No. 10, October, 2008
2071
without nitrite additives, as in region I. The initial oxidaꢀ
tion rate in the presence of nitrite ion with the leading
role of nitrogen dioxide according to the equilibrium
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2 H3O+ + NO2– + NO3
2 NO2 + 3 H2O
–
should be proportional to the square root of the nitrite
ion concentration, which is observed in the experiment.
The same mechanism of NO2ꢀmediated oxidation in the
initial step and in the acceleration step in the presence of
sodium nitrate in HNO at nH /nHNO = 10.17 is also
2O
3
3
indicated by similar temperature dependences of the rate
constants k1 and k2.
It may be assumed that with the further dilution of the
starting HNO3 solutions with water, the leading role in
the oxidation at the stage of the developed process will
transfer from NO2 to N2O3 and then to HNO2.
The presence of the carbonyl group increases the
reactivity of acetone compared to hydrocarbons. A comꢀ
parison of the oxidation rates of acetone and decane2
shows that the maximum rate of acetone oxidation at
6. A. I. Kazakov, D. A. Vaganov, L. P. Andrienko, E. Yu. Rubtsova,
Yu. I. Rubtsov, G. B. Manelis, Zh. Prikl. Khim., 2000, 73,
1270 [Russ. J. Appl. Chem., 2000, 73, 1342 (Engl. Transl.)].
7. O. S. Galyuk, Yu. I. Rubtsov, G. F. Malinovskaya, G. B. Manelis,
Zh. Fiz. Khim., 1965, 39, 2319 [Russ. J. Phys. Chem., 1965
(Engl. Transl.)].
92.8 °С and nH /nHNO = 2.92 is 350 times higher than
2O
3
that of decane oxidation.
8. J. Chedin, J. Chem. Phys., 1952, 49, 109.
Thus, the critical phenomena were observed in
the oxidation of acetone by nitric acid solutions. They
manifest themselves as an abrupt change in the apparent
rate constant for the initial oxidation step upon an insigꢀ
nificant change in the temperature or acid concentration.
The nature of these phenomena differs from that of the
known critical phenomena in the branched chain reacꢀ
tions of oxidation by molecular oxygen, which are related
to the dramatic change in the apparent rate constant in
the selfꢀacceleration step due to the difference in the rate
constants and orders of the chain branching and terminaꢀ
tion reactions. It seems important to study the nature of
the critical phenomena in the oxidation by bound oxygen.
It is assumed that they are due to changes in the structure
of an acid solution and, as a consequence, its solvating
ability with respect to the oxidizing agent or acetone molꢀ
ecule at a certain acid/water ratio in the solution.
9. N. M. Emanuelґ, D. G. Knorre, Kurs khimicheskoi kinetiki
[The Course of Chemical Kinetics], Vysshaya Shkola, Moscow,
1969, 431 pp. (in Russian).
10. A. I. Kazakov, L. P. Andrienko, Yu. I. Rubtsov, Zh. Fiz.
Khim., 1979, 53, 1054 [Russ. J. Phys. Chem., 1979 (Engl.
Transl.)].
11. G. V. Lagodzinskaya, N. G. Yunda, G. B. Manelis, Izv.
Akad. Nauk, Ser. Khim., 2006, 577 [Russ. Chem. Bull., Int.
Ed., 2006, 55, 597].
12. A. N. Nesmeyanov, N. A. Nesmeyanov, Nachala organicheskoi
khimii [The Elements of Organic Chemistry], Book 1, Khimiya,
Moscow, 1969, 663 pp. (in Russian).
13. D. Axente, Studia Universitatis BabesꢀBolyai. Physica, 2001,
Special Issue, 89.
14. A. P. Ballod, V. Ya. Shtern, Usp. Khim., 1976, 45, 1428
[Russ. Chem. Rev., 1976, 45 (Engl. Transl.)].
The authors are deeply grateful to A. N. Ivanova for
kindly provided program for the calculation of the kinetic
curves.
Received November 9, 2007;
in revised form March 25, 2008