Kinetics and mechanism of catalytic decomposition and oxidation of chlorine dioxide by the hypochlorite ion

Article abstract of DOI:10.1021/ic001106y

The oxidation of ClO₂ by OCl^-is first order with respect to both reactants in the neutral to alkaline pH range: -d[ClO₂]/dt = 2k_OCl[ClO₂][OCl^-]. The rate constant (T = 298 K, μ = 1.0 M NaCl

Full text of DOI:10.1021/ic001106y

Inorg. Chem. 2001, 40, 1833-1836
1833
Kinetics and Mechanism of Catalytic Decomposition and Oxidation of Chlorine Dioxide by
the Hypochlorite Ion
†
‡
,†
§
Viktor Csord a´ s, Bernie Bubnis, Istv a´ n F a´ bi a´ n,* and Gilbert Gordon
Department of Inorganic and Analytical Chemistry, University of Debrecen, P.O.B. 21,
Debrecen H-4010, Hungary, Novatek, A Division of EBB, Inc., Oxford, Ohio, and
Department of Chemistry and Biochemistry, Miami University, Oxford, Ohio
ReceiVed October 4, 2000
-
The oxidation of ClO2 by OCl is first order with respect to both reactants in the neutral to alkaline pH range:
-
-
d[ClO2]/dt ) 2kOCl[ClO2][OCl ]. The rate constant (T ) 298 K, µ ) 1.0 M NaClO4) and activation parameters
-
1
-1
q
q
are kOCl ) 0.91 ( 0.02 M s , ∆H ) 66.5 ( 0.9 kJ/mol, and ∆S ) -22.3 ( 2.9 J/(mol K). In alkaline
solution, pH > 9, the primary products of the reaction are the chlorite and chlorate ions and consumption of the
hypochlorite ion is not observed. The hypochlorite ion is consumed in increasing amounts, and the production of
the chlorite ion ceases when the pH is decreased. The stoichiometry is kinetically controlled, and the reactants/
products ratios are determined by the relative rates of the production and consumption of the chlorite ion in the
-
-
ClO2/OCl and HOCl/ClO2 reactions, respectively.
Introduction
implies that redox side reactions do not occur between chlorine
dioxide and other oxychlorine species. Deviations from this
general pattern were reported only in a few cases. The reaction
between the hypochlorite ion and chlorine dioxide was studied
The reactions between oxychlorine species are very sensitive
to the conditions applied and exhibit a very rich and complex
redox chemistry. As a consequence, chlorine may be stabilized
in different oxidation states, depending on the absolute and
relative reactant concentrations, pH, temperature, etc. Chlorine
dioxide quantitatively decomposes into the chlorite and chlorate
ions in alkaline solution, but it is reasonably stable in acidic
20
earlier, and it was proposed recently that a direct redox reaction
may occur between chlorine dioxide and the chlorite ion
-
18
coordinated to mercury(II) in the Hg(II)-ClO2 -ClO2 system.
The formation and subsequent decay of small amounts of
ClO2 were found in the reaction of HOCl with ClO2 at pH
-
1
2
aqueous solutions. The stability of the hypochlorite ion is also
pH-dependent. The overall third-order rate of decomposition
8
.82, but a detailed interpretation of these observations was not
2
1
-
given. It was concluded that the slow loss of chlorine dioxide
could be caused either by impurities or by a reaction with HOCl.
According to our preliminary studies, the transient concentration
of chlorine dioxide can be increased and systematically varied
by selecting appropriate experimental conditions. When OCl
is in excess over ClO2 , the decay of ClO2 is orders of
goes through a maximum at pH 7.10 (25 °C) and produces Cl
and ClO3- in a 2:1 ratio.3
Under acidic conditions, ClO2 was reported as a final product
-
4-9
in the HOCl-ClO2 reaction,
acid and metal ion catalyzed
-
1
0-19
decomposition of chlorite ion and related systems.
This
-
magnitude faster than the base-catalyzed disproportionation of
ClO2. In agreement with previous reports, this is a further piece
of evidence that HOCl either is involved in a direct redox
reaction with chlorine dioxide or somehow catalyzes its
decomposition.
*
To whom correspondence should be addressed.Phone: +36 52 512-
1
9
00, extension 2378. Fax: +36 52 489 667. E-mail: ifabian@delfin.klte.hu.
†
University of Debrecen.
Novatek, A Division of EBB, Inc.
Miami University.
‡
§
(
1) Gordon, G.; Kieffer, R. G.; Rosenblatt, D. H. Prog. Inorg. Chem. 1972,
5, 201.
2) The hypochlorite ion and hypochlorous acid are in fast protolytic
equilibrium, and their concentration ratio is determined by the pH. In
this paper, the two species will not be distinguished unless it is required
by the clarity of presentation.
1
Flis and co-workers reported the following rate law for the
(
2
0
HOCl-ClO2 reaction:
d[ClO2]
(
3) Adam, L. C.; F a´ bi a´ n, I.; Suzuki, G.; Gordon, G. Inorg. Chem. 1992,
-
) k [ClO ][HOCl] + k [ClO ]
(1)
I
2
II
2
31, 3534.
dt
(
(
(
(
(
(
4) Taube, H.; Dodgen, H. J. Am. Chem. Soc. 1949, 71, 3330.
5) Tang, T. F.; Gordon, G. EnViron. Sci. Technol. 1984, 18, 212.
6) Emmenegger, F.; Gordon, G. Inorg. Chem. 1967, 6, 633.
7) Aieta, E. M.; Roberts, P. V. EnViron. Sci. Technol. 1986, 20, 50.
8) Peintler, G.; Nagyp a´ l, I.; Epstein, I. R. J. Phys. Chem. 1990, 94, 2954.
9) Jia, Z.; Margerum, D. W.; Francisco, J. S. Inorg. Chem. 2000, 39,
This study was limited to relatively narrow concentration and
pH ranges, and the constraints of the applicability of the rate
law were not given. It was proposed that the second term in eq
1
corresponds to the decomposition of chlorine dioxide into Cl2
2614.
and O2. While this reaction is one of the possible fragmentation
(
(
10) Launer, H. F.; Tomimatsu, Y. J. Am. Chem. Soc. 1954, 76, 2591.
11) Launer, H. F.; Wilson, W. K.; Flynn, J. H. J. Res. Natl. Bur. Stand.
1
953, 51, 237.
(17) F a´ bi a´ n, I.; Gordon, G. Inorg. Chem. 1992, 31, 2144.
(18) F a´ bi a´ n, I.; S z_u cs, D.; Gordon, G. J. Phys. Chem. A 2000, 104, 8045.
(19) T o´ th, Z.; F a´ bi a´ n, I. Inorg. Chem. 2000, 39, 4608.
(20) Flis, I. E.; Mischenko, K. P.; Salnis, K. U. Russ. Appl. Chem. 1959,
32, 284.
(
(
(
(
(
12) Gordon, G.; Feldman, F. Inorg. Chem. 1964, 3, 1728.
13) Kieffer, R. G.; Gordon, G. Inorg. Chem. 1968, 7, 235.
14) Schmitz, G.; Rooze, H. Can. J. Chem. 1981, 59, 1177.
15) Schmitz, H.; Rooze, H. Can. J. Chem. 1985, 63, 975.
16) F a´ bi a´ n, I.; Gordon, G. Inorg. Chem. 1991, 30, 3994.
(21) Gordon, G.; Tachiyashiki, S. EnViron. Sci. Technol. 1991, 25, 468.
1
0.1021/ic001106y CCC: $20.00 © 2001 American Chemical Society
Published on Web 04/02/2001

1
834 Inorganic Chemistry, Vol. 40, No. 8, 2001
Csord a´ s et al.
[ClO3-]/
channels in condensed-phase photolysis of chlorine dioxide,22-29
it does not seem to be feasible in the thermally induced
decomposition. In support of this statement, it should be
remembered that slightly acidic solutions of chlorine dioxide
can be stored in the dark without apparent concentration loss
for several months.
In the present paper we report a pH- and temperature-
dependent kinetic study on the hypochlorite ion-chlorine
dioxide reaction. Our goal is to explore the exact kinetic role
of the hypochlorite ion in the decay of ClO2. The results are
also expected to contribute to a better understanding of the
Table 1. Stoichiometric Data for the Chlorine
Dioxide-Hypochlorite Ion Reactiona
4
∆[ClO-]/
[ClO
[ClO2-]/
[ClO
10 ×
pH
[ClO
2
]
0
(M)
2
]
0
2
]
0
2 0
[ClO ]
6
6.36
6.51
.34
9.20
9.20
3.10
3.10
9.03
9.03
3.10
9.20
9.20
9.20
9.20
5.78
5.78
3.10
5.78
9.20
9.20
5.78
5.78
9.20
3.10
3.10
0.44
0.44
0.04
0.05
0.03
0.00
-0.01
-0.01
0.00
0.01
0.01
0.07
0.06
0.08
0.29
0.30
0.47
0.48
0.48
0.50
0.50
0.49
0.49
0.50
0.93
0.92
0.95
0.98
0.96
0.96
0.99
0.98
0.98
0.91
0.93
0.93
0.70
0.66
0.52
0.50
0.49
0.50
0.49
0.49
0.49
0.49
6.94
7.25
7.26
7.35
0.46
0.45
7.43
7.45
0.49
0.50
0.45
0.44
0.48
0.23
-
mechanism of the HOCl-ClO2 reaction and other reactive
systems in which HOCl and ClO2 are present simultaneously.
7
7
7
8
.85
.85
.86
.21
Experimental Section
Chemicals. Sodium hypochlorite stock solutions were prepared as
8.32
8.69
8.98
8.99
3,8
0.01
-0.03
-0.03
0.02
-0.08
0.00
described earlier. Chlorine dioxide was generated by mixing aqueous
solutions of potassium persulfate and sodium chlorite. The product
ClO
steady stream of nitrogen gas. Stock solutions of NaOCl and ClO
stored at 5 °C in the dark and were standardized by iodometric titration
3
0
-
4
2
was passed into cold, dilute HClO
4
solution (∼10 M) with a
9.05
9.42
9.95
2
were
3
1
before use. To minimize loss to vaporization, ClO
were kept in and dispensed from a shrinking bottle.
2
stock solutions
1
1
0.23
0.49
1
9,32
All other
reagents were of reagent grade quality and used without further
purification. Samples were prepared using doubly deionized and
ultrafiltered or deionized and triply distilled water obtained from
MILLI-Q RG (Millipore) or Barnstead/NANOpur water purification
systems, respectively. With the exception of the temperature-dependent
studies, the kinetic measurements were made at 25 ( 0.1 °C and 1.0
a
-3
[
HOCl] ) 2.00 × 10 M. The concentrations of the individual
components were determined with an error less than (3%.
2 4
acidified to pH 2.0 with 0.5 M H SO . After addition of excess KI to
the sample, the iodine formed was titrated with standardized sodium
thiosulfate solution. Under such conditions, ClO
I while HOCl and ClO
The concentration of the unreacted hypochlorite ion was calculated by
combining the chromatographic and iodometric data. The concentration
of Cl was calculated from the mass balance for the chlorine species.
Ion chromatograms were recorded on a Dionex 100DX ion chro-
matograph equipped with AS9-HC analytical and AG9-HC guard
columns and using a conductivity detector. The eluent was 0.7 g/dm
Na CO . To avoid overloading of the analytical column, the ionic
2 3
strength was lowered to 0.1 M in these experiments.
Stock solutions of the reagents were analyzed by using the same
procedures. In these experiments, ClO
with nitrogen gas prior to analysis. The concentrations of the chlorite
and chlorate ions were less than 0.5% of the total chlorine in the ClO
and OCl solutions.
Time-resolved spectra were recorded on a HP-8543 diode array
spectrophotometer equipped with an Applied Photophysics RX-2000
-
3
does not react with
M ionic strength set with NaClO
Na CO . The pH was adjusted with phosphate or carbonate buffers.
Instrumentation and Methods. Iodometric titrations and pH
4
prepared from perchloric acid and
-
-
2
generate 2 and 4 equiv of iodine, respectively.
2
3
measurements were made with a Metrohm 721 NET Titrino system
equipped with Metrohm 6.0420.100 combined platinum and with
Metrohm 6.0202.000 combined glass electrodes, respectively. The pH
-
+
was measured within (0.003 unit, and it is defined as -log[H ] in
3
this paper.
2
In the stoichiometric experiments, a ClO stock solution was added
to a solution of HOCl and buffer in an all-glass, stoppered vessel. The
hypochlorite ion was used at least in 2 times excess in these
experiments. The reactor was tightly sealed immediately after introduc-
2
was purged from the samples
2
ing ClO . The experiments were designed such that the headspace above
the reaction mixture was always less than 1-2% of the total volume.
The reactants were thoroughly mixed with a Teflon stirring bar. The
spent reaction mixtures were analyzed typically 2 h after the reaction
was triggered. Within this time frame ClO was completely consumed
2
even in the slowest kinetic runs and subsequent reactions of the excess
HOCl were negligible.
2
-
3
5
rapid kinetics accessory. In agreement with a recent report, the light
intensity of the spectrophotometer was found to be strong enough to
initiate photochemical decomposition of ClO . Such a reaction was not
2
observed when the high-energy UV region of the light beam (<300
nm) was blocked out by inserting a polystyrene cutoff filter between
the light source and the sample. Kinetic traces at different wavelengths
An aliquot of the reaction mixture was diluted and treated with excess
-
33,34
ethylenediamine in order to remove OCl .
The sample was analyzed
-
-
and ClO
3
for ClO
2
by ion chromatography. Another aliquot was
-
were consistent with a simple first-order process when OCl was applied
(22) Mialocq, J. C.; Barat, F.; Gilles, L.; Hickel, B.; Lesigne, B. J. Phys.
in large excess over chlorine dioxide. The pseudo-first-order rate
constants, which were reproducible within 2%, were determined from
kinetic traces at 360 nm. The temperature dependence of the rate
constants was studied from 15 to 70 °C. Data fitting was made with
the program package SCIENTIST using nonlinear least-squares rou-
tines.36
Chem. 1973, 77, 742.
(
(
(
23) Dunn, R. C.; Anderson, J. L. J. Am. Chem. Soc. 1993, 115, 5307.
24) Chang, Y. J.; Simon, J. D. J. Phys. Chem. 1996, 100, 6406.
25) Pursell, C. J.; Conyers, J.; Denison, C. J. Phys. Chem. 1996, 100,
15450.
(
26) Thørgensen, J.; Jepsen, P. U.; Thomsen, C. L.; Poulsen, J. A.; Byberg,
J. R.; Keiding, S. R. J. Phys. Chem. 1997, 101, 3317.
27) Esposito, A. P.; Foster, C. E.; Beckman, R. A.; Reid, P. J. J. Phys.
Chem. 1997, 101, 5309.
28) Foster, C. E.; Reid, P. J. J. Phys. Chem. A 1998, 102, 3514.
29) Thørgensen, J.; Thomsen, C. L.; Poulsen, J. A.; Keiding, S. R. J. Phys.
Chem. A 1998, 102, 4186.
30) Rosenblatt, D. H.; Hayes, A. J. J. Org. Chem. 1963, 28, 2790.
31) Vogel, A. I. QuantitatiVe Inorganic Analysis, 3rd ed.; Longman:
London, 1961.
(
Results and Discussion
(
(
Stoichiometry. As shown in Table 1, a definite trend was
-
observed in the stoichiometry of the ClO2/OCl reaction as a
(
(
function of pH. At low pH, the data confirmed earlier results
-
-
and the main products were Cl and ClO3 :
(
(
32) Silverman, R. A.; Gordon, G. Anal. Chem. 1974, 46, 178.
33) Bolyard, M.; Fair, P. S.; Hautman, D. P. EnViron. Sci. Technol. 1992,
(35) Stanbury, D. M.; Figlar, J. N. Coord. Chem. ReV. 1999, 187, 223.
(36) SCIENTIST, version 2.0; Micromath Software: Salt Lake City, UT,
1995.
26, 1663.
(34) Adam, L. C.; Gordon, G. Anal. Chem. 1995, 67, 535.

Decomposition and Oxidation of ClO2
Inorganic Chemistry, Vol. 40, No. 8, 2001 1835
-
3
-
+
HOCl + 2ClO + H O ) 2 ClO + Cl + 3H
(2)
2
2
When the pH was increased, ClO2- was formed in increasing
-
amounts at the expense of Cl formation. At the same time,
-
-
the consumption of OCl and the [ClO3 ]produced/[ClO2]consumed
concentration ratio decreased. Above pH 9.0, the stoichiometry
corresponded to the disproportionation of chlorine dioxide and
could be approximated by the following equation:
-
-
2
-
2
ClO + 2OH ) ClO + ClO3 + H O
(3)
2
2
The stoichiometry could be given by the linear combination of
eqs 2 and 3 at any intermediate pH.
Figure 1. Experimental (b) and fitted on the basis of eq 5 (solid line)
-
pseudo-first-order rate constants, kobs, as a function of pH. [OCl ] )
-
3
3
.13 × 10 M, T ) 298 K, µ ) 1.0 M NaClO
4
.
While reaction 3 properly describes the stoichiometry of the
-
ClO2/OCl reaction in the alkaline pH range, it represents only
3
reported on the basis of an independent study before. The HOCl
path has a marginal contribution to the reaction, and the fitted
value for kOCl practically did not change when the second term
was neglected in eq 5.
-
-
an intermediate stage whenever OCl and ClO2 are present
simultaneously in the reaction mixture. The kinetics and
stoichiometry of the reaction between these two species have
been extensively studied.4 In alkaline solution it is a slow
-9
Temperature-dependent data at pH 8.0 were evaluated by
considering only the kOCl term in eq 5 and using earlier reported
thermodynamic data for the temperature dependence of Kp. The
activation parameters were calculated by fitting kOCl on the basis
of
-
-
process that produces ClO3 and Cl as final products. It follows
-
that the unreacted OCl would eventually remove the chlorite
ion produced in reaction 3 and the same stoichiometry would
prevail in the entire pH region.
Kinetics. The first-order decay of chlorine dioxide at large
excess of HOCl over ClO2 and the linear dependence of the
pseudo-first-order rate constant on the total concentration of the
hypochlorite ion confirm that the reaction is first order in both
reactants:
kT
h
e∆
Sq/R -∆Hq/(RT)
kOCl
)
e
(6)
q
q
The results are ∆H ) 66.5 ( 0.9 kJ/mol and ∆S ) -22.3 (
2
.9 J/(mol K).
Mechanism. The unique feature of the ClO2/OCl reaction
-
d[ClO2]
-
-
) k[ClO ][OCl ]
(4)
2
tot
is that the distinct pH dependence of the stoichiometry is not
associated with any change in the rate law as a function of pH.
The pseudo-first-order rate constant is strictly proportional to
dt
Thorough analysis of the kinetic traces gave no evidence of a
second term in the rate law as suggested by eq 1. This strongly
suggests that the first-order loss of chlorine dioxide reported
by Flis and co-workers20 was probably caused by some sort of
artifact in their experiments. At the high chlorine dioxide
concentrations used in that work, slow volatilization of ClO2
from the reaction vessel could possibly corrupt the data. It was
also confirmed in control experiments that the ClO2/OCl-
reaction was much faster than the decomposition of chlorine
dioxide with base over the entire pH region studied. Under
alkaline conditions (pH > 9.0), the hypochlorite ion was not
consumed (see eq 3) and the reaction could be formally treated
-
[OCl ] in the entire pH region, and the corresponding pH profile
exhibits an inflection point at the log Kp of HOCl (pH ) 7.40).
At the same time, the changeover between the limiting stoichi-
ometries (eqs 2 and 3) occurs in a narrow pH range at around
pH 8.1. These features imply that the same initial step is
operative regardless of the actual pH and that the stoichiometry
is controlled by the rates of subsequent competing reaction steps.
In accordance with these considerations, we propose the
following mechanism for the interpretation of the results:
-
+
OCl + H ) HOCl
fast equilibrium
(R7)
-
as an OCl -catalyzed decomposition of chlorine dioxide.
-
-
The pH dependence of kobs is shown in Figure 1. The
ClO + OCl ) Cl O
rate determining (R8)
2
2
3
characteristic sigmoid shape can readily be associated with the
-
-
-
acid-base equilibrium between HOCl and OCl . Apparently,
Cl O + ClO ) Cl O + ClO
(R9)
2
3
2
2
3
2
chlorine dioxide is oxidized much more quickly by the hy-
pochlorite ion than by the hypochlorous acid. Considering that
the acid-base reaction is a fast preequilibrium and taking into
account the stoichiometry of the reaction, kobs can be expressed
as follows:
-
-
Cl O + OH ) HOCl + ClO
3
(R10a)
2
3
or
-
+
Cl O + H O ) HOCl + ClO + H
(R10b)
+
2
3
2
3
kOCl
kHOClK [H ]
p
[OCl-]
kobs ) 2
+
(5)
tot
ClO2- + H ) HClO
+
(
+
+
)
fast equilibrium (R11)
1
+ K [H ] 1 + K [H ]
2
p
p
where Kp, kOCl, and kHOCl are the protonation constant of the
hypochlorite ion and the rate constants for the reactions of OCl
HOCl + HClO
2
) Cl
2
O
2
+ H
2
O
(R12)
(R13)
-
-
-
+
and HOCl, respectively.
Cl O + H O ) ClO + Cl + 2H
2
2
2
3
Fitting of kobs on the basis of eq 5 gave the following
-
1
-1
The formation of the Cl2O3- intermediate in the rate-
determining step is consistent with the simple second-order rate
law, and steps R7-R10 are sufficient to explain the stoichi-
results: log KP ) 7.40 ( 0.02, kOCl ) 0.91 ( 0.02 M s ,
-
3
-1 -1
and kHOCl ) (1.5 ( 0.3) × 10
M
s . The value for the
-
protonation constant of OCl is the same as the value that was

1836 Inorganic Chemistry, Vol. 40, No. 8, 2001
Csord a´ s et al.
ometry in alkaline solution. The reaction of ClO2 with HOCl,
if it occurs at all, is also expected to generate Cl2O3 via the
formation of HCl2O3. In this case, an immediate proton loss
should follow the rate-determining step.
The consumption of the chlorite ion is interpreted in terms
-
-
of the HOCl/ClO2 reaction (steps R11-R13). There seems to
be a consensus in the literature that this reaction proceeds via
-
the Cl2O2 intermediate and produces only ClO3 in excess
4
-9,21
A simple one-electron-transfer step between the reactants
would be an alternative of step R8, but this possibility can be
rejected on the basis of the following arguments. The primary
HOCl.
The mechanism postulated here assumes that the
-
stoichiometry of the ClO2/OCl reaction is determined by the
interplay of the two reaction paths leading to the formation and
consumption of the chlorite ion. The former path is relatively
fast in alkaline solution. This, coupled with the slow decay of
-
products of the electron transfer would be ClO2 and ClO.
According to earlier studies, ClO is a strong oxidant that would
-
-
immediately react with ClO2 (and excess ClO2) to form the
the chlorite ion, makes possible the accumulation of ClO2 .
chlorate and chloride ions.1 However, in this case the chlorite
ion would not be stable even at high pH and the mechanism
would contradict experimental observations. Furthermore, there
8,37
-
The rate of the initial step, R2, swiftly decreases as OCl is
converted to HOCl by decreasing the pH (step R1). In contrast,
-
+ 2
the rate of the HOCl/ClO2 reaction is proportional to [H ] in
the alkaline pH range and increases by about 6 orders of
-
is no reason to assume that the oxidation of ClO2 to ClO3 and
-
-
the reduction of OCl to Cl in less alkaline solution would
occur via respective lower and higher oxidation state intermedi-
ate chlorine species. It should be added that very recently it
was incorrectly proposed that ClO is capable of reducing ClO2
to the chlorite ion.38 Were this reaction step operative, it would
predict a very different stoichiometry than was observed
experimentally.
21
magnitude when the pH is decreased from 10.0 to 6.0. Thus,
the chlorite ion becomes a reactive intermediate at low pH and
-
it is completely consumed by the end of the ClO2/OCl reaction.
The opposite pH dependencies of the two kinetically coupled
reaction paths also explain the sharp changeover in the stoi-
chiometry as a function of pH.
In conlusion, the results presented here confirm that the
A relatively small negative entropy of activation was found
for the initial step. It is unlikely that solvation effects have
immediate products of the chlorine dioxide-hypochlorite ion
-
-
-
reaction are ClO3 and ClO2 . When OCl is used in excess,
the chlorite ion is involved in subsequent reactions and the
lifetime and transient concentration of this species are kinetically
controlled by the competing reaction paths. Perhaps the most
important implication of the results is that chlorine dioxide and
the hypochlorite ion cannot coexist in aqueous solution for an
extended period of time in the neutral to alkaline pH range,
and the reaction between these species may explain the enhanced
decomposition rate of chlorine dioxide in related systems. The
results also imply that under neutral conditions the chlorine
dioxide-hypochlorite ion reaction is competitive with the
formation of chlorine dioxide in the hypochlorous acid-chlorite
ion reaction. Thus, the corresponding mechanism must include
q
significant contribution to the activation parameters, and ∆S
is consistent with a weak association between the reactants in
the precursor complex. This interaction probably occurs between
-
the negatively charged oxygen atom of OCl and the partially
positive chlorine atom of ClO2, implying Cl-O-Cl connectivity
-
for the Cl2O3 intermediate.
According to the mechanism, the chlorite ion is produced in
-
the reaction of Cl2O3 with a second molecule of ClO2 (step
-
R9). This step is followed by the regeneration of OCl and the
-
formation of ClO3 . Step R10 is very fast, and with respect to
the rate law and overall stoichiometry, it is irrelevant whether
it occurs with a solvent molecule or with the hydroxide ion.
-
The two transient species Cl2O3 and Cl2O3 are presumably
-
the ClO2/OCl path for proper interpretation of the noted decay
formed in extremely low steady-state concentrations, and direct
experimental evidence is not available to confirm their existence.
However, our conclusions are corroborated by earlier kinetic
studies, which postulated the formation of very similar reactive
of ClO2 in that system.
Acknowledgment. This work was supported by NATO
Linkage Grant CRG.LG 973337 and by Grants OTKA T 029568
and M 028244 from the Hungarian National Research Founda-
tion.
1
,3,9,39-43
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6
Supporting Information Available: Two tables of observed rate
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constants as a function of OCl and ClO
2
concentrations, pH, and
temperature and a plot of observed rate constants as a function of
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[OCl ]. This material is available free of charge via the Internet at
http://pubs.acs.org.
IC001106Y