Acid Catalysis in Reaction of Ozone with Chloride Ions

Article abstract of DOI:10.1023/B:KICA.0000009047.90252.2d

The kinetics of the interaction of ozone with aqueous solutions of chlorides resulting in Cl₂ evolution to the gas phase was studied. The reaction of O₃ with Cl^- is accelerated by H ⁺ ions. The effects of the concentrations of H⁺ and Cl^-, the ionic strength, and temperature (ranged from 7 to 60°C) on the reaction rate were studied. A mechanism explaining the experimental kinetics was proposed. The acid catalysis is due to the formation of the HO ₃Cl complex, which is in equilibrium with H⁺, O ₃, and Cl^-. The constants of reactions involved in the proposed mechanism were determined.

Full text of DOI:10.1023/B:KICA.0000009047.90252.2d

Kinetics and Catalysis, Vol. 44, No. 6, 2003, pp. 740–746. Translated from Kinetika i Kataliz, Vol. 44, No. 6, 2003, pp. 810–816.
Original Russian Text Copyright © 2003 by Levanov, Kuskov, Zosimov, Antipenko, Lunin.
Acid Catalysis in Reaction of Ozone with Chloride Ions1
A. V. Levanov, I. V. Kuskov, A. V. Zosimov, E. E. Antipenko, and V. V. Lunin
Department of Chemistry, Moscow State University, Moscow, 119899 Russia
Received November 13, 2002
Abstract—The kinetics of the interaction of ozone with aqueous solutions of chlorides resulting in Cl2 evolu-
tion to the gas phase was studied. The reaction of O3 with Cl– is accelerated by H+ ions. The effects of the con-
centrations of H+ and Cl–, the ionic strength, and temperature (ranged from 7 to 60°C) on the reaction rate were
studied. A mechanism explaining the experimental kinetics was proposed. The acid catalysis is due to the for-
mation of the HO3Cl complex, which is in equilibrium with H+, O3, and Cl–. The constants of reactions involved
in the proposed mechanism were determined.
1
INTRODUCTION
ride, sodium perchlorate monohydrate of analytical purity
grade (Merck), and concentrated sulfuric acid of chemical
purity grade were used in the preparation of solutions.
Ozone was prepared by passing special purity oxy-
gen through a barrier discharge in an ozonizer 1. The
ozone concentration at the reactor inlet was measured
with a Medozon-254/5 photometer and normally was
10.0 g/m3. The flow rate of the initial gas mixture was
21 l/h in all experiments.
In the course of a run, we determined the concentra-
tion of Cl2 in the gas mixture at the reactor outlet and the
ozone concentration in the reaction solution; in some
cases we measured the pH of the reaction solution.
The composition of the gas mixture evolved was
analyzed on a Balzers PGM 407 mass spectrometer.
Along with the components of the initial mixture,
molecular Cl2 was found in the outgoing gas. Other gas-
eous reaction products were not found.
Studies of the interaction between ozone and chloride
ions in aqueous solutions are of current interest in the
light of the problem of changes in the concentration of
atmospheric ozone, as well as of the processes of water
preparation and purification. The reaction O3 + Cl– is
believed to be slow, and only rough estimates of its rate
constant at room temperature are known [1]. However,
our experimental findings showed that under certain
conditions the rate of this reaction can be considerably
increased. The aim of this work is to study the effect of
various experimental parameters (concentrations of H+
and Cl–, the ionic strength, and temperature) on the
reaction between ozone and chlorine ions in aqueous
solutions and its kinetic features.
EXPERIMENTAL
The concentration of Cl2 in the evolved gas mixture
was estimated by photometric iodometry. The gas mix-
ture passed through an oven (3) to decompose O3 and
then through a trap (4) with a solution of KI to absorb
Cl2 and to form the equivalent amount of I2, whose con-
centration was measured using a KFK-3 photometer.
Experiments were carried out using a setup whose
schematic is shown in Fig. 1. The reaction between
ozone and chlorine ions was performed in a tempera-
ture-controlled bubble-column reactor (2). The main
part of the reactor represented a glass cylinder (inner
diameter, 2.5 or 3 cm) with a sintered glass filter sealed
into its bottom for feeding the initial gas mixture.
Before the experiment, a solution under study was
placed in the reactor; the height of the liquid was 25–30 cm.
The exact volume of the solution in the reactor was cal-
culated from the previously measured density of solu-
tion and the weight in the reactor.
The amount of chlorine nCl entering the trap was cal-
2
culated as a function of time t from the known iodine
concentration and solution volume. The rate of the
O2
In most runs, solutions under study contained NaCl
and 0.01–1 M HCl, and [NaCl] + [HCl] = 1 M. In some
experiments, NaClO4 was added to vary the ionic
strength of solutions. An additional series of runs was
carried out with solutions containing 0.25 M HCl and
0–8 M H2SO4. Distilled water, standard solution of
hydrochloric acid (1 M), chemically pure sodium chlo-
1
3
2
4
Fig. 1. Schematic of the experimental setup: (1) ozonizer,
(2) bubble-column reactor, (3) oven for ozone decomposi-
tion, (4) trap with a KI solution.
1
Paper presented at the VI Russian Conference on Catalytic Reac-
tion Mechanisms (Moscow, October 1–5, 2002).
0023-1584/03/4406-0740$25.00 © 2003 MAIK “Nauka /Interperiodica”

ACID CATALYSIS IN REACTION OF OZONE WITH CHLORIDE IONS
741
[O3], µå
60
chlorine inflow in the trap dnCl /dt was estimated from
2
a tangent slope of the linear portion of the plot nCl vs. t.
2
The procedure for estimating Cl2 in the ozone-contain-
ing gas mixture was reported in detail in our previous
paper [2].
40
20
1
2
Estimation of the Concentration of Dissolved Ozone
To selectively determine the concentration of dis-
solved ozone in the presence of chlorine, we developed
a method based on the use of an alkaline solution of a
Cr(III) salt. In alkaline solutions ozone rapidly and
completely oxidizes Cr(III), which mainly exists in the
form of hydroxochromite [Cr(OH)6]3– [3], to chromate
0
1
2
3
Ionic strength, M
Fig. 2. Concentration of dissolved ozone vs. ionic strength
at room temperature (24°C). CO = 10 g/m . The composi-
3
CrO24– . This reaction was proposed for the determina-
3
tion of solution: (1) 0.5 M HCl + NaCl; (2) 0.5 M HCl +
tion of ozone in air [4]. We found that under these con-
ditions chlorine reacts slowly, and under certain condi-
tions one can neglect the formation of chromate via the
reaction of [Cr(OH)6]3– with ClO–. Chromate CrO24–
NaClO .
4
(ozone decay in solution is a first-order reaction), κ is
the rate coefficient of ozone dissolution, and CO is the
has a well-resolved absorption band with a maximum at
373 nm (the molar absorbance is 4610 l mol–1 cm–1) [5].
Therefore, the chromate concentration and hence the
ozone concentration can be measured by the absor-
bance at 373 nm using an alkaline Cr(III) solution with-
out ozone as a reference solution. To determine the
ozone concentration in the reaction solution, we used
the following procedure. A portion of the solution from
3
ozone concentration in the initial gas mixture. At room
temperature, the following estimates are made: κ ~
500 min–1 [6] and kO < 10 min–1, and, therefore,
3
Henry’s law [O3] = HCO is practically fulfilled.
3
The plot ln[O3] vs. 1/T is linear at temperatures 7–
the reactor was mixed with an equal volume of an alka- 41.5°C (see Fig. 3), since the concentration of dis-
line Cr(III) solution containing 0.005 M KCr(SO4)2 + solved ozone is determined by Henry’s law due to a rel-
([HCl] = 0.5 M) NaOH, where [HCl] is the HCl molar
concentration in the initial solution. Immediately after
mixing, the absorbance at 373 nm of the solution
formed was measured on a KFK-3 photometer. An
alkaline Cr(III) solution of the same composition
diluted with an equal volume of distilled water was
used as a reference solution. The ozone concentration
was determined from a calibration plot.
The concentrations of dissolved ozone found with
the above method are shown in Table 1 and Figs. 2 and
3. We found that under experimental conditions, the
concentration of dissolved ozone is determined by the
ionic strength only and is independent of the nature of
the supporting electrolyte (HCl, NaCl, or NaClO4)
(Fig. 2).
atively low value of kO . The concentration of dissolved
3
ozone at 60°C deviates from the linear plot in the
ln[O3] – 1/T coordinates because of a significant
increase in the kO value.
3
Table 1. Concentrations of dissolved ozone vs. the
[HCl]/[NaCl] ratio and temperature
[O3], µM
[HCl], M
7°C
24°C
41.5°C
60°C
0.01
0.1
0.3
0.4
0.5
0.7
1
–
–
27.3
–
11.5
11.9
11.1
–
73.1
–
42.3
44.9
–
Under our experimental conditions (a flow reactor
and a constant ozone concentration in the initial gas
mixture), dissolved ozone is present in a steady-state
concentration determined by the equality of the rates of
its dissolution and decay:
–
74.5
–
27.1
–
44.5
45.6
44.9
44
–
71.8
73.7
–
–
H
---------------------
[O3] =
CO ,
(1)
3
1 + k /κ
O3
–
–
Average value [O3] 73
27
11.5
where H is Henry’s constant (the ratio between the
equilibrium ozone concentrations in the solution and
3
Note: In the initial solution [HCl] + [NaCl] = 1 M and CO = 10 g/m .
gas phase), kO is the apparent rate constant of decay
3
3
KINETICS AND CATALYSIS Vol. 44 No. 6 2003

742
LEVANOV et al.
The experimental and simulation results showed
ln[O3] [µå]
that a nearly steady-state regime is quickly reached
when the concentrations of reactants in the reaction
solution, and the chlorine concentration in the outgoing
gas mixture remain almost constant. In the course of
runs (duration was at most 1.5 h), the reaction solution
was nearly of the same composition as the initial solu-
tion with the exception that ozone and chlorine
appeared. This result is due to the relatively slow
removal of chlorine from the reactor. Our flow system
did not operate in a strictly steady-state regime because
at least one of the reactants, chlorine ion, was not fed
during the run but was present within the reaction sys-
tem in a limited amount. In the case of a nearly steady-
state regime, the differential equation takes the follow-
ing form:
5
4
3
2
0.0028 0.0030 0.0032 0.0034 0.0036
1/T, 1/ä
Fig. 3. The natural logarithm of the average concentration of
the dissolved ozone vs. inverse temperature. CO = 10 g/m .
3
3
The composition of solution: [HCl] + [NaCl] = 1 M.
v
Vl
kapp[Cl–][O3] = CCl .
(4)
----
2
We found the heat of ozone dissolution to be
21 kJ/mol from the linear portion of the plot (Fig. 3),
and this value is in good agreement with previously
published data (17–24 kJ/mol) [7–9].
We used this expression for the calculation of the con-
stant kapp. The concentration of chlorine ions and the
volume of solution placed in the reactor are set by
experimentalist. The concentrations of ozone in the
solution and chlorine in the outgoing gas were mea-
sured with the above procedures. The rate of chlorine
Measurement of Apparent Rate Constant
of the Reaction O3 + Cl–
evolution from the reactor, vCCl , is equal to the exper-
2
imentally measured rate of the chlorine inflow to the
A flow-type bubbling reactor was used in our exper-
iments. In this case, the following differential equation
is applicable:
dnCl
2
-----------
trap: vCCl
=
. Hence, the final expression for the
2
dt
dSCl
calculation of kapp takes the form:
--------- = kapp[Cl–][O3]
dt
dnCl
(2)
1
2
----------- -----------------------------
kapp
=
.
(5)
–
v
Vl
[O3][Cl–]Vl
dt
----
– (kI + kII)[ClO ][O3] – CCl ,
2
where kapp is the apparent rate constant of the reaction
O3 + Cl–
ClO– + O2, kI and kII are the rate con-
Reaction Orders in Ozone and the Chloride Ion
stants of the reactions ClO– + O3
Cl– + 2O2 and
The expression for the calculation of kapp was found
assuming that the reaction has the first orders in ozone
and chlorine ion. To determine the reaction order, we
measured the rate of chlorine evolution from the reactor
based on the unit liquid volume in the reactor and deter-
ClO– + 2O3
ClO–3 + 2O2, respectively (values of kI
and kII are given in the work [10]), [X] is the concentra-
tion of substance X in a solution, CCl is the concentra-
2
tion of Cl2 in the gas phase (at the reactor outlet), SCl = mined its dependence on the concentration of ozone in
[ClO–] + [HOCl] + [Cl2] + [Cl–3 ] + CCl Vg/Vl is the
the initial gas mixture or on the concentration of chlo-
rine ions in the reaction solution (Fig. 4).
2
amount of the active chlorine in the reactor divided by
The relative rate of chlorine evolution is propor-
the volume of the liquid, v is the gas mixture flow rate
tional to the ozone concentration in the initial gas mix-
(in all experiments v = 21 l/h (standard conditions)), Vl
ture (see Fig. 4a) and, therefore, to the ozone concen-
is the volume of the liquid in the reactor, and Vg is the
tration in solution because, in our experiments at room
volume of the gas phase. At pH < 3, when almost all the
temperature, the ozone concentration in solution obeys
active chlorine is either molecular chlorine (Cl2) or Cl3–
Henry’s law.
ions in solution, the differential equation takes the form
dSCl
The relative rate of chlorine evolution is also pro-
portional to the concentration of chlorine ion in the
reaction solution (Fig. 4b). In experiments of this
series, we maintained a constant ionic strength and var-
ied the ratio of the concentrations of chlorine and per-
chlorate ions. The ozone concentration in the initial gas
v
–
----
--------- = kapp[Cl ][O3] – CCl ,
(3)
2
Vl
dt
where SCl = [Cl2] + [Cl–3 ] + CCl Vg/Vl.
2
KINETICS AND CATALYSIS Vol. 44 No. 6 2003

ACID CATALYSIS IN REACTION OF OZONE WITH CHLORIDE IONS
743
mixture was constant. Since the constant-temperature
dnCl2
dt Vl
24
1
µmol
l min
,
solubility of ozone in solutions used in our experiments
is only determined by the ionic strength (see above),
then in all the experiments of this series the ozone con-
centration in the reaction solutions was also constant.
Therefore, the experimental reaction orders with
respect to ozone and chlorine ions are unity.
(‡)
18
12
6
As shown below, the reaction under study is com-
plex. Hence, the term reaction order is barely applica-
ble to it. Nevertheless, the relative rate of chlorine evo-
lution is in fact nearly proportional to the concentration
of chlorine ions in the reaction solution.
0
10
20
30
40
[O3], g/m3
DISCUSSION
12
9
(b)
The apparent rate constant for the reaction O3 + Cl–
increases with increasing [H+] (Fig. 5). In highly acidic
media the rate of the reaction O3 + Cl– considerably
increases (Fig. 6). Hence, the H+ ions are the catalyst of
the reaction.
6
The apparent rate constant increases with an
increase in the solution ionic strength (Fig. 7) except
for the case when the ionic strength is created by NaCl.
An increase in the chloride concentration results in
some decrease in kapp (Fig. 7). A nonzero value of kapp
at [H+] = 0 (see Fig. 5) corresponds to an noncatalytic
reaction route.
3
0
0.25
0.50
0.75
1.00
[Cl–], å
Fig. 4. The rate of Cl evolution at room temperature (24°C)
2
These features can be rationalized assuming the fol-
lowing mechanism of the reaction between ozone and
chlorine ion in an aqueous solution:
divided by the volume of the liquid in the reactor vs. (a) the
ozone concentration in the initial gas mixture (composition
of solution: 0.01 M HCl + 0.9 M NaCl) and (b) the concen-
3
tration of chlorine ion in solution at CO = 10 g/m (compo-
3
Cl– + O3
ClO– + O2,
(I)
(II)
sition of solution: 0.1 M HCl, [NaCl] + [NaClO ] = 0.9 M).
4
H+ + O3 + Cl–
HO3Cl,
where k1 and k3 are the rate constants of steps (I) and
(III), respectively, and K2 is the equilibrium constant of
step (II). The quasi-equilibrium approximation for
HO3Cl implies that the concentration of dissolved
ozone in our experiments is the sum of the concentra-
tions of the O3 and HO3Cl species in solution. Accord-
ing to Eq. (6), the constant kapp is an increasing function
of [H+] and K2 for the nonnegative [H+] and [Cl–] values
and positive k1, K2, and k3 values (an additional condi-
tion should be [Cl–] < k3/k1) and is a decreasing function
of [Cl–]. Thus, expression (6) explains catalysis by
hydrogen ions and retardation by chlorine ions. The
effect of the ionic strength is due to an increase in the
concentration equilibrium constant K2 with an increase
in the ionic strength, while the rate constants k1 and k3
are independent of the ionic strength.
HO3Cl
HOCl + O2,
(III)
ClO– + Cl– + 2H+
HOCl + Cl– + H+
Cl2 + H2O or
Cl2 + H2O (fast).
(IV)
The hydration of the species is not shown here for sim-
plicity. According to the proposed mechanism, the
reaction occurs via two parallel pathways. One of them,
noncatalyzed reaction (I) includes the formation of
hypochlorite ion in the reaction of chlorine ion with
ozone.
The catalyzed reaction pathway includes the forma-
tion of the intermediate HO3Cl from H+, O3, and Cl– in
a fast equilibrium step (II) followed by HO3Cl decom-
position to HOCl in step (III). Protonation of ClO– and
HOCl in acidic media due to the fast equilibration steps
(IV) produces Cl2. The expression for the apparent rate
constant under quasi-equilibrium approximation for
HO3Cl takes the form
The effect of H+ on the rate of ozone reactions in
acidic media can be due to the formation of protonated
ozone HO+3 , which is more reactive than normal ozone
[11–15]. In our case, the assumption of the formation of
k1 + k3K2[H+]
HO+3 can explain catalysis by the H+ ion but cannot
----------------------------------------
kapp
=
,
(6)
1 + K2[H+][Cl–]
explain the retardation by chlorine ions. Meanwhile,
KINETICS AND CATALYSIS Vol. 44 No. 6 2003

744
LEVANOV et al.
k
1.00
app, l mol–1 min–1
k
app, l mol–1 min–1
0.3
0.2
0.1
7°C
24°C
0.75
0.50
0.25
0
0.2
0.4
0.6
0.8
1.0
0
0.2
0.4
0.6
0.8
[H+], å
[H+], å
4
3
2
1
12
9
41.5°C
60°C
6
3
0
0.2
0.4
0.6
0.8
0
0.1
0.2
0.3
0.4
[H+], å
[H+], å
–
+
Fig. 5. The apparent rate constant of the reaction O + Cl vs. [H ] at various temperatures. The composition of solution: [HCl] +
3
[NaCl] = 1 M.
k
app, l mol–1 min–1
k
app, l mol–1 min–1
9
6
3
1.2
0.8
0.4
0
1
2
3
0
2
4
6
8
[H2SO4], å
1
2
3
Ionic strength, M
–
Fig. 6. The apparent rate constant of the reaction O + Cl
3
–
vs. concentration of sulfuric acid in solution at room tem-
perature (24°C). The composition of solution: 0.25 M HCl +
H SO . The concentration of dissolved ozone was evalu-
Fig. 7. The apparent rate constant of the reaction O + Cl
3
vs. ionic strength at room temperature (24°C). The compo-
sition of solution: 0.5 M HCl, and either: (1) NaCl,
2
4
ated on the basis of data presented in [11].
(2) NaClO , or (3) KNO .
4
3
Based on experimental plots kapp vs. [H+] (Fig. 5)
the assumption for the formation of HO3Cl gives a rea-
sonable explanation for all the experimental features of
for solutions with the ionic strength of 1 M, we found
the kinetics of the reaction between O3 and Cl– in aque- the numerical values of the constants k1, K2, and k3 and
ous solutions. The intermediate HO3Cl can be treated as the characteristics of their temperature dependences
a complex of protonated ozone with chlorine ion.
(Table 2).
KINETICS AND CATALYSIS Vol. 44 No. 6 2003

ACID CATALYSIS IN REACTION OF OZONE WITH CHLORIDE IONS
745
lnk1 [l mol–1 min–1]
lnK2 [(l/mol)2]
2
3
1
2
3
1
0
–1
–2
–3
–5
0.0028
–4
0.0038 0.0028
1/T, 1/ä
0.0033
0.0033
0.0038
1/T, 1/ä
lnke [min–1]
ln(ke K2) [l2 mol–2 min–1]
4
5
3
1
3
2
–1
–3
0.0028
1
0.0030
0.0034
0.0038
1/T, 1/ä
0.0033
0.0038
1/T, 1/ä
Fig. 8. The natural logarithm of the constants k , K , k , and k K vs. inverse temperature: (1) our data (ionic strength 1 M); (2) data
1
2
3
3 2
of [16] (ionic strength, 0.3 M); 3, the upper limit of k estimated in [1] (ionic strength, 1–4 M).
1
The experimental plot kapp vs. [H+] (Fig. 5) is linear proposed mechanism) at room temperature was
at 7°C and, therefore, it is possible to estimate the prod- found, which is also in good agreement with our data
uct of constants k3K2 = 0.21 l2 mol–2 min–1 and the upper (Fig. 8).
limit of K2 (or the lower limit of k3). The temperature
The value k1 = 0.12 l mol–1 min–1 (room tempera-
dependences of the constants k1, K2, k3, and k3K2 are
ture, ionic strength 1 M) found in [17] practically coin-
presented in Fig. 8.
cides with the value presented in this work (see Table 2).
The activation energy or the enthalpy of reaction
Thus, the findings of our work give the most accu-
rate description of the kinetics and mechanism of inter-
action between ozone and chlorine ion in aqueous solu-
tions and more adequately characterize the contribution
of this interaction in the processes of atmospheric
chemistry.
steps are as follows:
Step (I)
79
Step (II)
55
Step (III)
18
Our experimental findings are in agreement with
published data. For instance, data on the catalysis of the
reaction of ozone with chlorine ion by H+ ion were pre-
sented in [16], according to which the apparent rate
constant linearly depends on [H+] and is independent
of [Cl–]. The authors of [16] only partly characterized
the effect of [H+] on kapp because they only performed
experiments at low temperatures and in a narrow
range of H+ concentrations: 0°C, [H+] = 0.01–0.2 M;
9.5°C, [H+] = 0.02–0.1 M. The parameters of the linear
plot kapp vs. [H+] found in [16] (k1 and k3K2 values in
terms of the mechanism proposed by us) are in good
agreement with our data.
Table 2. Parameters of the reaction steps
k1,
k2,
k3,
T, °C
l mol–1 min–1
(l/mol)2
min–1
7
0.021
0.13
0.71
4.7
<0.05
0.10
0.32
1.3
>4
11
17
24
24
41.5
60
Note: Ionic strength, 1 M. The composition of the solution: [HCl] +
[NaCl] = 1 M.
In the work [1], the upper limit of the rate con-
stant of the noncatalyzed pathway (k1 in terms of the
KINETICS AND CATALYSIS Vol. 44 No. 6 2003

746
LEVANOV et al.
9. Kvanaugh, M.C. and Trussel, R.R., J. Am. Water Works
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