Oxyhalogen-Sulfur Chemistry: Oxidation of Hydroxymethanesulfinic Acid by Chlorite

Article abstract of DOI:10.1021/jp953795i

The reaction between chlorite and hydroxymethanesulfinic acid, HOCH2SO2H, (HMSA) has been studied in the pH range 4-8.The reaction is very complex and has a variable stoichiometry that is dependent on ratio of the oxidant to the reductant.In excess HMSA, the stoichiometry is ClO2(1-) + HOCH2SO2H -> SO4(2-) + HCHO + Cl(1-) + 2H(1+) while in excess ClO2(1-) the stoichiometry is 3ClO2(1-) + 2HOCH2SO2H -> 2SO4(2-) + 2HCOOH + 4H(1+) + 3Cl(1-).The reaction is very fast and is characterized by a short induction period (ca. 1 s) which is followed by a rapid and autocatalytic ClO2 production.The first stage in the oxidation is the cleavage of C-S bond to form SO4(2-) and HCHO.In the presence of excess oxidant the HCHO is oxidized to HCOOH.After prolonged standing (72 h or more), the HCOOH can be oxidized to CO2; giving a third possible stoichiometry.The rate-determining step is a 2-electron oxidation of HMSA by ClO2(1-) to give the sulfonic acid HOCH2SO3H.A 19-reaction mechanism is used to simulate the reaction.There is reasonable agreement between the experiments and simulations.

Full text of DOI:10.1021/jp953795i

J. Phys. Chem. 1996, 100, 9377-9384
9377
Oxyhalogen-Sulfur Chemistry: Oxidation of Hydroxymethanesulfinic Acid by Chlorite1
Mohamed A. Salem
Department of Chemistry, UniVersity of Qatar, P.O. Box 2713 Doha, Qatar
Cordelia R. Chinake and Reuben H. Simoyi*
Department of Chemistry and Center for Nonlinear Science, West Virginia UniVersity,
Morgantown, West Virginia 26506-6045
ReceiVed: December 21, 1995; In Final Form: March 21, 1996X
The reaction between chlorite and hydroxymethanesulfinic acid, HOCH2SO2H, (HMSA) has been studied in
the pH range 4-8. The reaction is very complex and has a variable stoichiometry that is dependent on ratio
-
2-
of the oxidant to the reductant. In excess HMSA, the stoichiometry is ClO2 + HOCH2SO2H f SO4
+
+
HCHO + Cl- + 2H+ while in excess ClO2 the stoichiometry is 3ClO2 + 2HOCH2SO2H f 2SO4
-
-
2-
2HCOOH + 4H+ + 3Cl-. The reaction is very fast and is characterized by a short induction period (≈1 s)
which is followed by a rapid and autocatalytic ClO2 production. The first stage in the oxidation is the cleavage
of C-S bond to form SO42- and HCHO. In the presence of excess oxidant the HCHO is oxidized to HCOOH.
After prolonged standing (72 h or more), the HCOOH can be oxidized to CO2; giving a third possible
-
stoichiometry. The rate-determining step is a 2-electron oxidation of HMSA by ClO2 to give the sulfonic
acid HOCH2SO3H. A 19-reaction mechanism is used to simulate the reaction. There is reasonable agreement
between the experiments and simulations.
Introduction
mate22 and chlorite23 it shows clock reaction characteristics. The
sulfur center passes through many oxidation state intermediates
Very little progress has been made over the years in studies
of kinetics and mechanisms of reactions which involve sulfur
compounds. There has been no real need for such studies until
quite recently when it was discovered that sulfur compounds
can participate in a variety of reactions that give rise to complex
dynamical behavior.2 Complex dynamics such as chemical
oscillations,3 pattern formation,4 and chemical chaos5 had
previously been observed only in connection with oxyhalogen
chemistry.6-8 The chemical reactivity of some sulfur amino
acids in living organisms has not yet been fully understood
because of inadequate knowledge on the kinetics and mecha-
nisms of the sulfur end of the amino acid.9 The prevalence of
sulfur compounds as environmental pollutants has also increased
our need to know more about sulfur chemistry in general.10
The knowledge gathered so far on the chemistry of sulfur
compounds highlights its complexity. Mechanistic studies on
sulfur chemistry are hampered by existence of free radical
mechanisms,11 autoxidations,12 and general propensity of sulfur
compounds to aggregate.13
The complex reaction behavior of sulfur compounds has given
rise to exotic kinetics such as oligooscillations,14 chemical
oscillations,15 chemical wave propagation,16 and the novel pH-
driven oscillations.17 A better understanding of this complex
behavior will be possible after kinetics and mechanism of
relevant reactions have been characterized.
The chlorite ion exhibits exotic nonlinear dynamics in nearly
all its reactions.18 Chlorite is the most prolific oxyhalogen
species in production of clock reactions and chemical oscilla-
tors.19 The ion has its own built-in nonlinearity in its reactions
in the form of HOCl autocatalysis.20
on its way to being oxidized to sulfate (oxidation state change
-2 to +6).24 There have been numerous speculations as to the
type of intermediates formed. The presently accepted sequence
is a series of two-electron oxidation processes which produce
progressively the sulfenyl, sulfinic, and sulfonic acids.25 By
using one of the postulated intermediates as the starting material,
the oxidation mechanism of thiourea can be better evaluated.
In this paper we report on the kinetics and mechanism of
oxidation of hydroxymethanesulfinic acid (HMSA) by chlorite
in acidic media. HMSA is a stable sulfinic acid (in basic
environments) in which the sulfur center is at an oxidation state
of +2. Oxidation of a sulfur center in excess oxidant can take
it to the +6 oxidation state (SO42-), a process accompanied by
a cleavage of the C-S bond in HMSA. Using a compound
with a +2 sulfur oxidation state eliminates many other possible
intermediates (compared to starting with sulfur at oxidation state
-2) and reduces the possibility of polymerizations.
The motivation for this study is to attempt to systematically
evaluate kinetics and mechanism of sulfur compound reactions
which are of interest in nonlinear chemistry. Reactions of
chlorite with thiourea26 and formamidinesulfinic acid (FSA)
yield traveling chemical27 wavefronts of ClO2, as well as
oscillatory behavior in a CSTR. A complete explanation for
such behavior is also not possible until the kinetics and
mechanism of the reactions involved are known. This study
of the kinetics and mechanism of the chlorite-HMSA reaction
is thus part of our continuing efforts at elucidating the kinetics
and mechanisms of reactions involving sulfur compounds.1
A secondary motivation for our study is the importance of
HMSA in industrial applications and atmospheric chemistry.
Under the trade names of Rongalite and Aldanil, HMSA is a
major constituent of dyes28 and bleaches. In the troposphere,
HMSA is known to reversibly decompose to formaldehyde and
SO2(aq) (or HSO3-) in fog and cloud water, thus contributing
Thiourea is one of the sulfur compounds which exhibits
various nonlinear features when reacted with several oxidizing
agents.21 With mild oxidizing agents such as iodine, it shows
autoinhibitory and oligooscillatory behavior14 while with bro-
X Abstract published in AdVance ACS Abstracts, May 1, 1996.
quite significantly to the S(IV) balance in the atmosphere.29
A
S0022-3654(95)03795-6 CCC: $12.00 © 1996 American Chemical Society

9378 J. Phys. Chem., Vol. 100, No. 22, 1996
Salem et al.
-
mechanistic study of HMSA’s oxidation can help in improving
dye properties as well as sulfur abatement procedures.
The excess ClO2 is used for the further oxidation of
formaldehyde:38
HCHO + H2O f HCOOH + 2H+ + 2e-
(R3)
Experimental Section
Materials. The following reagent-grade chemicals were used
without purification: hydroxymethanesulfinic acid monosodium
salt dihydrate, (Aldrich), perchloric acid, 72% (Fisher), and
sodium perchlorate (Fisher). All solutions were prepared using
singly-distilled water. Technical grade sodium chlorite (Aldrich)
was 78-81% pure. The sodium chlorite was recrystallized once
from a water-methanol mixture to bring assay value to 96%.
The recrystallized chlorite was standardized iodometrically by
adding acidified potassium iodide and titrating liberated iodine
against sodium thiosulfate with freshly prepared starch as
indicator.30 Chlorine dioxide was prepared by a standard
method of oxidizing potassium chlorate in acidic medium.31
Chlorine dioxide was stored in acidic medium in volumetric
flasks wrapped in aluminum foil at 4 °C. The chlorine dioxide
was then standardized by its molar absorptivity coefficient of
1265 M-1 cm-1 at 360 nm32 on a Perkin-Elmer Lambda 2S
UV/vis spectrophotometer.
For ratios between 1 and 1.5 the product solution contained
a mixture of formaldehyde and formic acid. Sulfate formation
in stoichiometry R2 was quantitative based upon the initial
HMSA concentrations. In high excess ClO2-, R > 3 and after
prolonged standing (24 h or longer), formic acid is further
oxidized to CO2:39
HCOOH f CO2 + 2H+ + 2e-
(R4)
Reaction R4 is so slow that a clean stoichiometric determi-
-
nation of the ClO2 - HMSA reaction could not be deduced
even after 48 h. The decomposition of ClO2- 40 occurs on same
time scale as reaction R4, and thus after prolonged standing,
the ClO2- oxidant concentration decreased as ClO2 accumulated.
Qualitative tests during and at end of the reaction were
negative for sulfite.41 Thus the postulated reversible decom-
position of HMSA into sulfite and HCHO42 does not seem to
-
Methods. All experiments were carried out at 25 °C and at
an ionic strength of 0.50M (NaClO4). The ClO2-/HMSA
reaction was monitored spectrophotometrically at λ ) 360 nm.
Kinetics measurements were performed on a Hi-Tech Scientific
SF-61AF stopped-flow spectrophotometer. The data from the
spectrophotometer were amplified and digitized via an Omega
Engineering DAS-50/1 16-bit A/D board interfaced to a Tandon
386SX computer for storage and data analysis. Stoichiometric
determinations were perfomed by analyzing for ClO2-, HCHO,
occur at a rate compatible with the time scale of the ClO2
-
HMSA reaction.
-
Reaction Kinetics. The ClO2 - HMSA reaction presents
reaction dynamics which are much more complex than those
normally obtained from oxyhalogen-sulfur chemistry.1 In the
UV/visible region the reaction mixture has no absorbing species
except for ClO2 at 360 nm. This was confirmed by a three-
dimensional reaction kinetics profile which confirms lack of
spectral activity except for the peak at 360 nm which grows
with time.
HCOOH, SO4 , Cl-, and HMSA. Sulfate was analyzed
2-
gravimetrically as BaSO4
while HMSA35 and ClO2 were
33,34
-
The reaction itself was very fast, often going to completion
in 2 s or less. The reactions with R > 1 were all characterized
by a short induction period (0.2 to 1 s) in which no activity
was observed in the absorbance at 360 nm. The induction
period gave way to a rapid formation of ClO2 (Figure 1a).
Formation of ClO2 is in three parts. In the first stage there is
a gentle increase in rate of formation of ClO2. In the second
stage the rate starts to decrease but then rapidly picks up again
in an autocatalytic manner, heading to an abrupt halt. In the
third stage the ClO2 continues to increase at a very slow rate
that is not on the same time scale as the first two stages. The
induction period is inversely proportional to initial chlorite
concentration [ClO2-]0 (Figure 1b). The maximum chlorine
dioxide absorbance, after the second stage, (λ ) 360 nm) shows
a direct relationship with [ClO2-]0 with saturation (Figure 1c).
The reaction had a complex dependence on [HMSA]0 (Figure
2a). At higher [HMSA]0, with fixed [ClO2-]0, formation of
ClO2 only shows a single autocatalytic step. Maximum ClO2
absorbance also shows a positive dependence on [HMSA]0, and
at constant [ClO2-]0, the induction period decreases with
[HMSA]0 (Figure 2b). This was an unexpected result as most
clock reactions are expected to show retardation in rate as the
concentration of the reductant is increased at constant oxidant
concentrations.
analyzed titrimetrically.30 Spot tests were used for HCHO,35
2-
HCOOH,36 and Cl-. For quantitative determination of SO4
,
reaction solutions were allowed to stand for at least 24 h before
addition of BaCl2. The precipitate was also allowed to settle
for 24 h before filtering, drying, and weighing.
Reactions were perfomed in excess chlorite so as to utilize
the formation of chlorine dioxide as an indicator reaction.
HMSA is known to be a powerful reducing agent especially in
the presence of alkali and is also known to decompose to carbon
monoxide, carbon dioxide, and oxides of sulfur in presence of
dilute acids.37
Results
The stoichiometry of the reaction is complex and variable.
At least three plausible stoichiometries could be deduced from
varying initial conditions. In approximately equimolar and
excess HMSA conditions the stoichiometry was determined as
ClO2- + HOCH2SO2H f
SO42- + HCHO + Cl- + 2H+ (R1)
Qualitative analysis tests confirmed the aldehyde product
(Tollen’s reagent).35 The complete consumption of ClO2- was
confirmed by the inability of the product solution to oxidize
acidified iodide.30
The reaction also had a complex dependence on pH (Figure
3a). The instability of HMSA at low pH37 did not permit the
running of reactions below pH 4. In general, the reaction is
faster in lower pH conditions, although no relationship was
evident. At constant [ClO2-]0 and [HMSA]0, the maximum
ClO2 concentrations did not seem to vary with pH, except for
the expected decrease at higher pH due to ClO2 hydrolysis.43
The traces in Figure 3a were obtained by running the reaction
in different buffers. Predictable behavior is obtained in condi-
tions of pH 4.0. The reaction dynamics change completely
-
In excess ClO2 (3 > R ) [ClO2-]/[HMSA] > 1), the
stoichiometry became
3ClO2- + 2HOCH2SO2H f
2SO42- + 2HCOOH + 4H+ + 3Cl- (R2)

Oxyhalogen-Sulfur Chemistry
J. Phys. Chem., Vol. 100, No. 22, 1996 9379
Figure 2. (a, top) Absorbance traces at 360 nm showing the effect of
-
[HMSA]0 in slightly alkaline conditions. [ClO2
] ) 0.0015 M;
0
[HMSA]0 ) (a) 0.000 15, (b) 0.000 25, (c) 0.000 35, (d) 0.000 55, and
(e) 0.000 65 M. (b, bottom) Effect of [HMSA]0 on the induction period
for data in Figure 2a. ClO2 is not formed in stoichiometric excess of
HMSA.
when pH goes below 4. The reaction is also self-destructing.
A reaction run in an unbuffered solution at pH 5 will have a
pH of approximately 3 at the end of the reaction as the reaction
produces H+ as an important product (see stoichiometries R1,
R2, and R3). The limit of the reaction’s reproducibility can be
seen in Figure 3b. By varying acid concentrations gradually,
the reaction profile can be maintained for some range of [H+]0.
At a critical [H+]0 (typically ≈ pH 4), the reaction profile
changes dramatically as the system becomes essentially irre-
producible. In Figure 3b, reproducibility is obtained for traces
a-d, and irreprodicibility becomes evident in traces e and f as
the acid is increased. Direct reaction between HMSA and ClO2
was also monitored. The reaction is very fast (Figure 4) in both
excess HMSA and excess ClO2.
Figure 1. (a, top) Absorbance traces at 360 nm at varying initial
chlorite concentrations and in slightly alkaline conditions (pH ) 7.3).
[HMSA]0 ) 0.0005 M; [ClO2-]0 ) (a) 0.001, (b) 0.002, (c) 0.003, (d)
0.004, and (e) 0.005 M. The arrows on trace (b) indicate the end of the
first period and the end of the second period. (b, middle) Relationship
between [ClO2-]0 and the induction period for data in Figure 1a. (c,
bottom) Effect of [ClO2-]0 on the maximum [ClO2] obtained after the
end of the second period. Maximum absorbance goes to zero when
[ClO2-]0 ) [HMSA]0 ) 0.0005 M. ClO2 obtained only when ClO2- is
in stoichiometric excess.
The rate of the reaction between ClO2- and HCHO was also
estimated. Upon mixing the two reagents (neutral to slightly
acidic pH) there was an immediate formation of ClO2 after a
very short induction period (Figure 5). The rapid formation of
-
ClO2 indicates either that the oxidation of HCHO by ClO2 is

9380 J. Phys. Chem., Vol. 100, No. 22, 1996
Salem et al.
Figure 4. Rapid reaction between ClO2 and HMSA. The reaction is
faster than the ClO2-/HMSA reaction and is catalyzed by ClO2-. [ClO2]0
-
) 0.000 77 M; [HMSA]0 ) 0.0004 M. Trace (a) no ClO2-; (b) [ClO2
) 0.000 02 M.
]
0
Figure 3. (a, top) Effect of pH on the reaction. HMSA rapidly
decomposes at pH lower than its pKa. [ClO2-]0 ) 0.003 M; [HMSA]0
) 0.0005 M. pH ) (a) 4.0, (b) 4.3, (c) 7.0, (d) 7.3, (e) 7.6, and (f) 7.8.
(b, bottom) Effect of acid over a small range of [H+]0. As [H+]0
increases, the reaction becomes irreproducible at some critical [H+]0
which is determined by the initial reactant concentrations. [ClO2-]0 )
0.003 M; [HMSA]0 ) 0.002 M. [H+]0 ) (a) 6.0 × 10-5, (b) 7.0 ×
10-5, (c) 8.0 × 10-5, (d) 9.0 × 10-5, (e) 1.0 × 10-4, and (f) 1.5 ×
10-4 M.
Figure 5. Oxidation of HCHO by ClO2- in slightly acidic pH. Rapid
accumulation of ClO2 indicates a rapid ClO2-/HCHO reaction. [ClO2-
) (a) 0.001, (b) 0.002, (c) 0.003, and (d) 0.004 M. [HCHO]0 ) 0.00133
M.
]
0
The first step in this oxidation is the 2-electron oxidation of
HMSA by the Cl(III) species:
very fast, or that the ClO2 formed reacts very slowly with
HCHO. It appears the answer is a combination of both since
the reaction of ClO2 with HCHO was observed to be very slow.
The ClO2/HCHO reaction has a half-life of more than 1 h which
is ineffective on our reactions’ time scale. Thus the rapid
buildup of ClO2 in the ClO2-/HCHO reaction is due to the very
slow ClO2 consumption rate by HCHO and HCOOH.
HOCH2SO2H + ClO2- + H+ f
HOCH2SO3H + HOCl (R5)
Two possible routes exist for sulfonic acid: one involves
2-
further oxidation to SO4
:
HOCH2SO3H + H2O f SO42- + HCHO + H+ +2e- (R6)
Mechanism
and another is an initial hydrolyis of the sulfonic acid followed
by an oxidation:
The experimental data suggest that the complete oxidation
of HMSA is the prerequisite for ClO2 formation. The observa-
tion that the reaction between ClO2 and HMSA is complete
within a fraction of a second (Figure 4), while that between
ClO2- and HMSA requires about 0.5 s to get to the end of the
induction period appears to confirm this.
HOCH2SO3H + H2O f HSO4- + CH3OH + H+ (R7)
CH3OH f HCHO + 2H+ + 2e-
(R8)

Oxyhalogen-Sulfur Chemistry
J. Phys. Chem., Vol. 100, No. 22, 1996 9381
TABLE 1: Mechanism of the Oxidation of
Hydroxymethanesulfinic Acid by Chlorite
Paths R6 and (R7 + R8) are kinetically indistinguishable if
R5 is the rate-determining step.
reaction
In excess oxidant, the organic residue, HCHO, can be further
oxidized to formic acid:
no.
reaction
M1 ClO2- + H+ a HClO2
M2 2ClO2 + Cl- + H2O f HClO2 + ClO2- + HOCl
M3 ClO2- + HOCl + H+ a Cl2O2 + H2O
HCHO + H2O f HCOOH + 2H+ + 2e-
(R9)
M4 Cl2O2 + 2ClO2- + 2H+ a 2ClO2 + 2HOCl
M5 ClO2- + HOCH2SO2H + H+ f HOCl + HOCH2SO3H
M6 ClO2- + HOCH2SO3H f HOCl + SO42- + H2CdO + H+
M7 ClO2- + H2CdO + H+ f HOCl + HCOOH
M8 HOCl + HOCH2SO2H f Cl- + HOCH2SO3H + H+
M9 HOCl + HOCH2SO3H f Cl- + SO42- + H2CdO + 3H+
M10 HOCl + H2CdO f Cl- + HCOOH + H+
M11 HOCl + HCOOH f CO2 + Cl- + H2O + H+
M12 2ClO2 + HOCH2SO2H + H2O f
Any of the oxychlorine species can act as the oxidant. Since
reaction R9 is of same order in rate as the ClO2--HMSA
reaction, the coupling of these two reactions will produce very
complex overall kinetics.
2ClO2- + HOCH2SO3H + 2H+
A further reaction is oxidation of HCOOH:
M13 2ClO2 + HOCH2SO3H + H2O f 2ClO2- + H2CdO + SO4
M14 2ClO2 + H2CdO + H2O f HCOOH + 2ClO2- + 2H+
M15 Cl2O2 + HOCH2SO2H + H2O f 2HOCl + HOCH2SO3H
M16 Cl2O2 + HOCH2SO3H + H2O f
2-
HCOOH f CO2 + 2H+ + 2e-
(R10)
2HOCl + SO42- + H2CdO + 2H+
M17 Cl2O2 + H2CdO + H2O f 2HOCl + HCOOH
M18 HOCl + Cl- + H+ f Cl2 + H2O
Reaction R10, being very slow, will not affect the overall
reaction dynamics on the time scale of the HMSA oxidation
reaction.
M19 Cl2 + 2ClO2- f 2ClO2 + 2Cl-
rapidly produce Cl2 which can undergo reactions analogous to
R12, R13-type reactions:
-
ClO2 has to be oxidized to form ClO2, yet in the reaction
being studied, ClO2- is reduced by the sulfur compound to Cl-.
ClO2 is formed only when the ClO2- is in stoichiometric excess
HOCl + Cl- + H+ f Cl2(aq) + H2O
2ClO2- + Cl2 f 2ClO2 + 2Cl-
ClO2- + Cl2 f Cl2O2 + Cl-
(R15)
(R16)
(R17)
-
over reductant. In the presence of excess oxidant the ClO2
can disproportionate after prolonged sitting in acidic solution:
44
5ClO2- + 4H+ f 4ClO2 + Cl- + 2H2O
(R11)
The sequence R15-R17 may be kinetically inconsequential
if R16 and R17 are of the same (or higher) order of magnitude
in rate as R12 and R13.
Induction Period. During this period no ClO2 is formed.
Either there is not enough HOCl to make reaction R12 effective,
or that all ClO2 formed is quickly reduced by the reducing
species in solution. ClO2 is thermodynamically unstable with
Reaction R11, however, is too slow to be effective in 1-5 s.
The formation of ClO2 is very rapid and appears autocatalytic.
The HOCl intermediate plays a crucial role in accelerating rate
of production of ClO2. The generally accepted ClO2--HOCl
reaction has the stoichiometry:45
-
respect to ClO2 in the presence of a reducing agent (either a
one- or a two-electron reductant):
2ClO2- + HOCl + H+ f 2ClO2 + Cl- + H2O
(R12)
-
ClO2 + e- f ClO2
(R18)
In general reaction R18 is slow especially with a 2-electron
reductant and normally starts slowly as ClO2- accumulates, and
one observes a sigmoidal decay curve.46 Further reduction of
The stoichiometry of the ClO2 formation is ultimately reaction
R11 but catalyzed by HOCl. Autocatalysis in reaction R11 can
be explained by the asymmetric intermediate Cl2O2, which yields
2 HOCl molecules when subjected to a 2-electron reduction:
ClO2 yields Cl- which acts to produce more ClO2
:
-
-
2ClO2 + Cl- + H2O f 2ClO2- + HOCl + H+
(R19)
ClO2- + HOCl + H+ f Cl2O2 + H2O
Cl2O2 + 2H+ + 2e- f 2HOCl
(R13)
(R14)
The reaction mixture has enough Cl- to make R19 important
and render the rate of ClO2 consumption rapid. Thus in this
system the ClO2--HMSA reaction is fast enough such that the
appearance of ClO2 indicates total consumption of HMSA.
Kinetic data, however, indicates that the ClO2--HCHO and
ClO2-HCOOH reactions are not as rapid, and thus concurrent
accumulation of ClO2, HCHO, and HCOOH is possible within
the reaction mixture.
Sequence R13 + R14 shows normal quadratic autocatalysis
in HOCl.38 After reaction R5, which acts as an initiator, most
of the oxidation of substrate will be done by HOCl. The
autocatalytic HOCl production thus can explain the observed
-
nonlinear kinetics. While HOCl can react with ClO2 as in
Computer Simulations
R13, it can also simultaneously react with the other substrates,
especially organic residues. Possible reductants in solution
include HOCH2SO2H, HOCH2SO3H, CH3OH, HCHO, HCOOH,
and Cl-. With the exception of HCOOH and Cl- all the
reductants are expected to react rapidly and irreversibly with
HOCl to produce Cl-. The reaction of HOCl with Cl- should
A simplified mechanism suitable for computer simulation was
distilled from the comprehensive mechanism described above.
Table 1 shows our proposed 19-reaction scheme. The scheme
only has two types of reactions: normal oxychlorine reactions
and oxychlorine-sulfur reactions. Standard sulfur-sulfur

9382 J. Phys. Chem., Vol. 100, No. 22, 1996
Salem et al.
TABLE 2: Rate Constants and Rate Laws Used in the Numerical Simulations
reaction no.
V (forward rate)
V (reverse rate)
M1
M2
M3
M4
M5
M6
M7
M8
1 × 109[ClO2-][H+]
3.16 × 107[HClO2]
5 × 102[ClO2][ClO2][Cl-]
5 × 106[ClO2-][HOCl][H+]
6.6 × 102[Cl2O2]
9.8 × 105[Cl2O2][ClO2-][ClO2-][H+][H+]
2.1 × 104[ClO2-][HOCH2SO2H][H+]
1.0 × 103[ClO2-][HOCH2SO3H]
4 × 103[ClO2-][H2CdO][H+]
1.5 × 105[HOCl][HOCH2SO2H]
2.5 × 103[HOCl][HOCH2SO3H]
6 × 104[HOCl][H2CdO]
1 × 10-2[ClO2][ClO2][HOCl][HOCl]
M9
M10
M11
M12
M13
M14
M15
M16
M17
M18
M19
1 × 10-1[HOCl][HCOOH]
6 × 104[ClO2][ClO2][HOCH2SO2H]
1 × 10-1[ClO2][ClO2][HOCH2SO3H]
3 × 10-2[ClO2][ClO2][H2CdO]
2.4 × 104[Cl2O2][HOCH2SO2H]
8 × 103[Cl2O2][HOCH2SO3H]
6 × 105[Cl2O2][H2CdO]
1.8 × 104[HOCl][Cl-][H+]
1.1 × 101[Cl2]
1.8 × 106[Cl2][ClO2-][ClO2
]
-
interactions were ignored since there was no experimental
evidence for any polymerizations of the substrate. The existence
of only one intermediate, HOCH2SO3H, before formation of
product, SO42-, justifies this assumption.
but only very slowly with ClO2. Reactions M7, M10, M14,
and M17 involve the oxidation of HCHO. On the time scale
of ClO2--HMSA reaction, it is doubtful reaction R7 would
contribute significantly to the mechanism, which is why the
oxidation of CH3OH was not considered. While oxidation of
HCHO is facile, the oxidation of HCOOH is not. Oxidation of
HCOOH requires a high acid environment and a catalyst.39 The
oxidation of HCOOH will be the slowest step in this mechanism.
(c) Formation of ClO2: In the presence of HOCl, formation
The 19-reaction scheme is quite exhaustive, and it was arrived
at after considering all the possible reactions in solution and
discarding those that would have been ineffective based on rate
considerations and abundance of reactive species. There are
four oxidizing species in solution, ClO2, ClO2-, HOCl, and Cl2,
and four reducing species, HOCH2SO2H, HOCH2SO3H, HCHO,
and HCOOH (excluding CH3OH, whose existence is specula-
tive). A permutation of these species should give 16 possible
reactions. Control experiments of HCOOH with ClO2 and
-
of ClO2 from ClO2 is very rapid. This reaction (M3 + M4)
is on the same time scale as reactions M12-M14. Accumula-
tion of ClO2 will depend on the balance between M3 + M4
and reactions that consume ClO2.
-
-
ClO2 showed no activity on the time scale of the ClO2
-
(d) Autocatalysis: The sharp increase in [ClO2] at end of the
induction period has been rationalized in this mechanism by
the quadratic HOCl autocatalysis evident in M3 + M4.
Intermediate Cl2O2 is then involved in the oxidation of all the
reducing agents in solution: reactions M4 and M15-M17.
The mechanism in Table 1 was simulated using semiimplicit
Runge-Kutta methods implemented for stiff systems of ordinary
differential equations by Kaps and Rentrop.48 Table 2 shows
the rate laws and rate constants used in the simulations. Only
the oxychlorine reactions were made reversible (M1, M3, M4,
and M18). The kinetics constants used for these 4 reactions as
well as reactions M2 and M19 were available from literature.
Reactions M5-M17 involve an oxidation by an oxychlorine
species of a sulfur compound or an organic residue. These were
assumed to be irreversible, thus reducing the number of kinetics
variables needed.
Estimation of unknown rate constants: The rate constant for
reaction M5, kM5, was estimated from this study using induction
period data (e.g., Figure 1a). A very good estimate of the rate
constant could be made by assuming that the total consumption
of HMSA is a prerequisite for the end of the induction period.
The initial estimate for kM6 was derived from an earlier study
involving the reaction of chlorite and thiourea.26 kM7 was
estimated from this study (see data shown in Figure 5). The
value of kM8 was not very significant in the simulations as long
as it was higher than kM5. The value accepted for kM9, 2.5 ×
103 M-1 s-1, was determined by the best fit for the simulations.
The initial estimate for kM11 was taken from similar work done
by Shilov et al.49 kM12 was also estimated from the data in Figure
1.
HMSA reaction and were not considered.
Cl2 is a powerful oxidizing agent, but it has mostly been
omitted in the mechanism. This is because of the expected
quantity of Cl2 (very low) and the reversibility of reaction M18.
Reaction M19 should be able to halt the accumulation of Cl2.
Addition of three more reactions involving the oxidation by
chlorine of HOCH2SO2H, HOCH2SO3H, and HCHO did not
alter the observed reaction dynamics in the simulations although
it made the reaction system much stiffer.47 Cl2, if formed, is
unstable with respect to Cl- and ClO2 in presence of excess
-
ClO2 (see sequence R15-17).
Although reactions involving oxidation by ClO2 are very
-
slow compared to the HOCl oxidation reactions, ClO2- reactions
have, however, been used in this mechanism because ClO2- is
the most abundant oxyhalogen species and would contribute
significantly to the rate of reaction on the basis of pure mass-
action kinetics (M3, M5-M7).
The most important features of the mechanism in Table 1
are the following:
(a) Consumption of HMSA: This is initiated in reaction M5
by a 2-electron oxidation of HMSA to give sulfonic acid and
HOCl. HOCl is a much more powerful oxidizing agent and
will quickly react with more reductant to give Cl- (reaction
M8). The rate-determining step for the ClO2--HMSA reaction
is thus reaction M5. It is known that sulfonic acids are very
stable and are not easy to further oxidize. Reactions M6, M9,
M13 and M16 involve irreversible cleavage of the C-S bond
2-
to form SO4
.
(b) Oxidation of the organic residues: Reactions M6, M9,
M13, and M16, apart from forming SO42-, also form the organic
The simulations were most sensitive to kinetics parameters
for reactions M3, M5, M9, M12, and M13. The simulations
correctly predicted the induction period and the quantity of ClO2
residue, formaldehyde, HCHO. Control experiments have
-
shown that HCHO reacts rapidly with HOCl, Cl2, and ClO2
,

Oxyhalogen-Sulfur Chemistry
J. Phys. Chem., Vol. 100, No. 22, 1996 9383
not oligooscillatory, though ClO2 production is very complex.
The most important reaction in the mechanism is M5 (from the
table) which initiates the reaction and forms HOCl. The balance
between reactions M12-M14 and reaction M4 determines the
global dynamics of the reaction. The importance of M4 reaction
over M12-M14 ensures nonlinear kinetics as the autocatalytic
pathway contributes more to the observed global reaction
dymanics. The rapid HCHO-HOCl/Cl2 reaction coupled with
the relatively slow HCHO-ClO2 reaction voids the possibility
for oligooscillations. The end of the induction period gives way
to autocatalytic ClO2 production, hence the sigmoidal ClO2
production curve.
-
Reaction R11 gives the stoichiometry of the ClO2 dispro-
portionation in acidic medium. The availability of reactive
intermediates, HOCl, Cl2 from reduction of ClO2- fuels a rapid
ClO2 production. This explains the variable rate of ClO2
-
production. ClO2 first oxidizes the organic residue:
ClO2- + HCHO + H+ f HCOOH + HOCl (R20)
The HOCl produced can also oxidize HCHO:
HCHO + HOCl f HCOOH + Cl-
(R21)
or participate in reactions R13 and R14 to autocatalytically
produce ClO2. Thus rate of ClO2 production will be propor-
tional to the rate of formation of HOCl in R19. The slower
ClO2-HCHO reaction permits the autocatalytic buildup of ClO2.
After the total oxidation of HCHO, the production of ClO2
decreases as no further formation of HOCl takes place. The
very slow ClO2 production at the end of the reaction is due to
the slow HCOOH oxidation:
ClO2- + HCOOH + H+ f CO2(g) + HOCl + H2O (R22)
which produces the reactive intermediate, HOCl, at a very slow
-
rate. The reactive intermediate then reacts with ClO2 (R12)
to produce ClO2. This explains the very slow ClO2 production
at the end of the reaction. The proposed mechanism has also
been able to explain Figure 2b in which the induction period
decreased with [HMSA]0. Reaction R5 initially produces HOCl
which then takes over the bulk of the oxidations. The HOCl
also produces ClO2 (reaction R12) which can be rapidly
consumed only as long as HMSA still exists in solution.
Figure 6. (a, top) Comparison between experimental (solid line) and
calculated (circles) results at conditions: [HMSA]0 ) 0.0005 M,
[ClO2-]0 ) 0.003 M. The reaction is not buffered. Species followed is
ClO2. (b, bottom) Simulations results for other species in solution that
cannot be followed spectrophotometrically. Solid line [HCHO], dotted
line [SO42-]; squares [HMSA]; filled circles [HOCH2SO2H].
HOCH2SO3H, HCHO, and HCOOH all react very slowly with
-
ClO2. Higher [HMSA]0 rapidly produces HOCl from ClO2
,
thus fuelling the reaction and diminishing the induction period.
formed (Figure 6a). The simulations also show that the sulfonic
acid, HOCH2SO3H, rises to a maximum before falling at the
end of the induction period (Figure 6b). This could not be
observed experimentally but was conjectured to be happening.
Of interest in Figure 6b is the time dependence of [HCHO]
(solid line). The aldehyde is formed during the course of the
reaction, reaching a peak at the point where [HMSA] goes to
zero. It then decays, vanishing at the point where SO42- attains
its (stoichiometric) maximum concentration value. The simula-
tions show that HCHO can coexist with ClO2 since the rate of
HCHO oxidation by ClO2 is slow (from Figure 6a,b).
Acknowledgment. We are grateful for helpful discussions
with Prof. S. B. Jonnalagadda. This work was supported by
the West Virginia EPSCoR program. We thank the University
of Qatar for granting leave of absence to M.A.S.
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