S-oxygenation of thiocarbamides I: Oxidation of phenylthiourea by chlorite in acidic media

Article abstract of DOI:10.1021/jp0458654

The oxidation of 1-phenyl-2-thiourea (PTU) by chlorite was studied in aqueous acidic media. The reaction is extremely complex with reaction dynamics strongly influenced by the pH of reaction medium. In excess chlorite concentrations the reaction stoichiometry involves the complete desulfurization of PTU to yield a urea residue and sulfate: 2ClO₂ + PhN(H)CSNH ₂ + H₂O → SO₄^2- + PhN(H)CONH₂ + 2Cl^- + 2H⁺. In excess PTU, mixtures of sulfinic and sulfonic acids are formed. The reaction was followed spectrophotometrically by observing the formation of chlorine dioxide which is formed from the reaction of the reactive intermediate HOCl and chlorite: 2ClO₂^- + HOCl + H⁺ → 2ClO₂(aq) + Cl^- + H₂O. The complexity of the ClO₂ ^- - PTU reaction arises from the fact that the reaction of ClO ₂ with PTU is slow enough to allow the accumulation of ClO ₂ in the presence of PTU. Hence the formation of ClO₂ was observed to be oligooscillatory with transient formation of ClO₂ even in conditions of excess oxidant. The reaction showed complex acid dependence with acid catalysis in pH conditions higher than pK_a of HClO ₂ and acid retardation in pH conditions of less than 2.0. The rate of oxidation of PTU was given by -d[PTU]/dt = k₁[ClO₂ ^-][PTU] + k₂[HClO₂][PTU] with the rate law: -d[PTU]/d/ = [Cl(III)]τ[PTU]₀/K_a1 + [H + ] [K ₁K_a1 + k₂[H + ]]; where [Cl(III)]_T is the sum of chlorite and chlorous acid and K_a1 is the acid dissociation constant for chlorous acid. The following bimolecular rate constants were evaluated; k₁= 31.5±2.3 M^-1 s ^-1 and k₂ = 114±7 M^-1 s^-1. The direct reaction of ClO₂ with PTU was autocatalytic in low acid concentrations with a stoichiometric ratio of 8:5; 8ClO₂ + 5PhN(H)CSNH₂ + 9H₂O - 5SO₄^2- + 5PhN(H)CONH₂ + 8Cl^- + 18H⁺. The proposed mechanism implicates HOCl as a major intermediate whose autocatalytic production determined the observed global dynamics of the reaction. A comprehensive 29-reaction scheme is evoked to describe the complex reaction dynamics.

Full text of DOI:10.1021/jp0458654

J. Phys. Chem. A 2005, 109, 1081-1093
1081
S-Oxygenation of Thiocarbamides I: Oxidation of Phenylthiourea by Chlorite in Acidic
Media
Tabitha R. Chigwada,† Edward Chikwana, and Reuben H. Simoyi*
Department of Chemistry, Portland State UniVersity, Portland, Oregon 97207-0751
ReceiVed: September 13, 2004; In Final Form: NoVember 19, 2004
The oxidation of 1-phenyl-2-thiourea (PTU) by chlorite was studied in aqueous acidic media. The reaction is
extremely complex with reaction dynamics strongly influenced by the pH of reaction medium. In excess
chlorite concentrations the reaction stoichiometry involves the complete desulfurization of PTU to yield a
-
2-
urea residue and sulfate: 2ClO2 + PhN(H)CSNH2 + H2O f SO4 + PhN(H)CONH2 + 2Cl- + 2H+. In
excess PTU, mixtures of sulfinic and sulfonic acids are formed. The reaction was followed spectrophoto-
metrically by observing the formation of chlorine dioxide which is formed from the reaction of the reactive
intermediate HOCl and chlorite: 2ClO2- + HOCl + H+ f 2ClO2(aq) + Cl- + H2O. The complexity of the
-
ClO2 - PTU reaction arises from the fact that the reaction of ClO2 with PTU is slow enough to allow the
accumulation of ClO2 in the presence of PTU. Hence the formation of ClO2 was observed to be oligooscillatory
with transient formation of ClO2 even in conditions of excess oxidant. The reaction showed complex acid
dependence with acid catalysis in pH conditions higher than pKa of HClO2 and acid retardation in pH conditions
-
of less than 2.0. The rate of oxidation of PTU was given by -d[PTU]/dt ) k1[ClO2 ][PTU] + k2[HClO2]-
[PTU] with the rate law: -d[PTU]/dt ) [Cl(III)]T[PTU]0/Ka1 + [H + ] [k1Ka1 + k2[H + ]]; where [Cl(III)]T is
the sum of chlorite and chlorous acid and Ka1 is the acid dissociation constant for chlorous acid. The following
bimolecular rate constants were evaluated; k1 ) 31.5(2.3 M-1 s-1 and k2 ) 114(7 M-1 s-1. The direct
reaction of ClO2 with PTU was autocatalytic in low acid concentrations with a stoichiometric ratio of 8:5;
8ClO2 + 5PhN(H)CSNH2 + 9H2O f 5SO42- + 5PhN(H)CONH2 + 8Cl- + 18H+. The proposed mechanism
implicates HOCl as a major intermediate whose autocatalytic production determined the observed global
dynamics of the reaction. A comprehensive 29-reaction scheme is evoked to describe the complex reaction
dynamics.
Introduction
It was reported in a 1961 publication that phenylthiourea is
extremely toxic to rats but that diphenylthiourea was not.11 There
Sulfur compounds undergo a variety of metabolic reactions,
namely oxidations, reductions, hydrolysis and conjugations.1
Sulfur in most organic configurations is nucleophilic, and
nucleophilic atoms are usually susceptible to metabolic oxida-
tion.2 The oxidation of sulfur-containing compounds represents
an important aspect of sulfur metabolism. Oxidations of sulfur
compounds appear to be involved in many cellular functions,
which include the reductive degradation of polypeptide hor-
mones and proteins, regulation of protein synthesis, maintenance
of intracellular redox potential, and protection of the cell from
oxidative damage.3
The molecular basis for S-oxygenation of sulfur compounds
is not well understood since relatively few studies have
addressed this problem. Although S-oxygenation of organosulfur
xenobiotics by microsomes supplemented with NADPH and
oxygen has been extensively studied, the facile oxidations of
sulfur compounds by hydrogen peroxide and other species of
reduced oxygen, such as superoxide ion, have not received the
same attention. In general, the metabolism of sulfur-based
xenobiotics such as thionicotinamide4 and thiobenzamide5-7 is
mediated by enzyme catalysis, with both the microsomal
cytochrome P-450 system and the flavin-containing monooxy-
genases8-10 playing a role.
was and has never been a justification why such physiological
differences exist in very similar compounds. No doubt the
metabolic pathways of the two compounds must be different,
but how, and why?
Subtle differences in substituted thioureas can impart vastly
different degrees of toxicity. For example, while R-naphthyl-
thiourea (ANTU) produces pulmonary edema and pleural
effusion in rats, its oxidation gives R-naphthylurea, an innocuous
substance to rats.12 Toxicity and tumorigenicity of R-naph-
thylthiourea can then be evaluated from these experimental data.
The possibilities are limited to either the parent compound,
metabolic intermediates, or the sulfur-based leaving groups. The
toxicity can be narrowed down to only these few processes and
compounds.
We report, in this manuscript, on the oxidation of phenylth-
iourea by chlorite in acidic medium. The aim of this study was
to evaluate the mechanism, products, and intermediates formed
in this oxidation. The S-oxygenation mechanism of phenyl-
* Corresponding author.
† Permanent address: Department of Chemistry, West Virginia Univer-
sity, Morgantown, WV 26506-6045.
10.1021/jp0458654 CCC: $30.25 © 2005 American Chemical Society
Published on Web 01/26/2005

1082 J. Phys. Chem. A, Vol. 109, No. 6, 2005
Chigwada et al.
thiourea by chlorite would contribute greatly to the possible
elucidation of its metabolic activation. Phenylthiourea contains
the thiocarbamide group that is responsible for goitrogenic
activity9,13 in methimazole14 and 6-propyl-2-thiouracil.15 Its use
as a goitrogenic was discontinued due to its strong toxicological
side effects.16,17 Of special interest is the fate of the C-S bond
after the full oxidation of PTU. Its cleavage, or lack of, would
give us a better insight into the speculation of a possible
mechanism for its toxicity.
Experimental Section
Materials. The following analytical grade chemicals were
used without any further purification: perchloric acid (72%),
1-phenyl-2-thiourea, PTU (Acros), sodium perchlorate (Fisher),
and barium chloride (Aldrich). Commercially available sodium
chlorite (Aldrich) varied in purity (78-88%) with the main
impurities being chloride and carbonate. The sodium chlorite
was recrystallized once from a water/ethanol/acetone mixture
to bring the assay value to approximately 96%. The recrystal-
lized chlorite was standardized iodometrically by adding excess
acidified potassium iodide and titrating liberated iodine against
sodium thiosulfate with freshly prepared starch as an indicator.18
Chlorine dioxide was prepared by the standard method of
oxidizing potassium chlorate in acidic medium.19 The aqueous
chlorine dioxide was stored in an acidic medium in a volumetric
flask wrapped in aluminum foil at 4EC. The chlorine dioxide
was then standardized iodometrically and results obtained were
also confirmed spectrophotometrically by using a molar ab-
sorptivity coefficient of 1265 M-1 cm-1 at 360 nm on a Perkin-
Elmer Lambda 2S UV/Vis spectrophotometer.20
Tests for AdVentitious Metal Ion Catalysis. Recent work by
Doona et al. showed adventitious metal ions catalysis in
reductions involving thiourea.21 It was important to eliminate
the possibility for such catalysis in the chlorite-PTU system.
Water used for preparing reagent solutions was obtained from
a Barnstead Sybron Corporation water purification unit capable
of producing both distilled and deionized water (Nanopure). Not
much difference was observed in the general reaction kinetics
observed with distilled and deionized water. We utilized
inductively-coupled plasma mass spectrometry (ICPMS) to
quantitatively evaluate the concentrations of a number of metal
ions in the water used for our reaction medium. ICPMS analysis
showed negligible concentrations of iron, copper, and silver and
approximately 1.5 ppb of cadmium and 0.43 ppb in lead.
Reactions run with metal ion chelators such as EDTA and
deferroxamine did not offer different kinetics and dynamics.
Methods. All experiments were carried out at 25(0.5 °C
(Neslab thermostat bath) and an ionic strength of 1.0 M (sodium
perchlorate). The ClO2-, ClO2/PTU reactions were monitored
spectrophotometrically at λ ) 360 nm as reactions were
performed in excess chlorite so as to utilize the formation of
chlorine dioxide as an indicator of the extent of the reaction.
These reactions were also concurrently monitored at 249 nm,
which is the absorption peak for phenylthiourea. Figure 1a shows
superimposed UV spectra of the three major reagents in the
reaction mixture: phenylthiourea (λmax ) 249 nm, ꢀmax ) 12 312
M-1 cm-1), phenylurea (λmax ) 235 nm, ꢀmax ) 37 180 M-1
cm-1), and chlorine dioxide (λmax ) 360 nm, ꢀmax ) 1265 M-1
cm-1). The absorptivity coefficient of chlorine dioxide at 249
nm was 130 M-1 cm-1 while both phenylthiourea and phenyl-
urea had negligible contributions to the absorbances observed
at 360 nm. Kinetics measurements were performed on a Hi-
Tech Scientific SF61-DX2 double-mixing stopped-flow spec-
trophotometer. The data from the spectrophotometer were
Figure 1. (A) Three superimposed spectra of (a) 5 × 10-5
M
phenylthiourea (reductant), (b) 1 × 10-5 M phenylurea (product), and
(c) 8.8 × 10-4 M chlorine dioxide (product) showing peaks at 248,
235, and 360 nm, respectively. Phenylurea has a reasonably high
absorptivity coefficient at 249 nm (λmax for phenylthiourea). At the
beginning of the reaction, however, all absorbance activity observed
at 249 nm will be due to the depletion of phenylthiourea. There is no
interference from substrate or product at 360 nm. (B) Spectral scans
of reagent solutions containing phenylthiourea and varying concentra-
tions of chlorite between 200 and 350 nm. PTU has a well-defined
peak at 249 nm which progressively diminishes with increasing chlorite
concentrations. At chlorite-to-phenylthiourea ratios greater than 2.0 a
chlorine dioxide peak is observed at 360 nm after a short induction
period. [PTU]0 ) 5.0 × 10-4 M, [H+]0 ) 0.125 M, [ClO2-]0 ) (a) 2.5
× 10-5 M, (b) 5.0 × 10-5 M, (c) 1.0 × 10-4 M, (d) 2.0 × 10-4 M, and
(e) 3.5 × 10-4 M.
amplified and digitized via an Omega Engineering DAS-50/1
16-bit A/D board interfaced to a Pentium IV computer for
storage and data analysis. Stoichiometric determinations were
carried out by mixing various ratios of chlorite and PTU in
stoppered volumetric flasks and scanning them spectrophoto-
metrically for formation of ClO2 and depletion of PTU after an
incubation period of up to 2 days. Qualitative and quantitative
analysis of sulfate was performed through its precipitation as
BaSO4. For reactions run in excess chlorite conditions, the
excess chlorite oxidizing power was evaluated by addition of
excess acidified iodide which was titrated against standard
thiosulfate with freshly prepared starch as indicator. The
thiosulfate titer was next plotted against the amount of chlorite
used and this plot was extrapolated to zero titer to obtain the
stoichiometry of the reaction.

S-Oxygenation of Thiocarbamides: I
J. Phys. Chem. A, Vol. 109, No. 6, 2005 1083
Synthesis of Phenylthiourea Dioxide (N-Phenylaminoimi-
nomethanesulfinic Acid).22 Exactly 0.02 mol of PTU with 0.07
mol of molybdenyl acetylacetonate was dissolved in a 40-60
mixture of ether and dioxane and was carefully and slowly added
to 2 equiv of hydrogen peroxide with continuous stirring at -15
°C. This mixture was allowed to stand for an hour at room
temperature and then left overnight at -4 °C. After thawing,
some off-white powdery crystals were obtained which could
be washed with acetonitrile. These phenylthiourea dioxide
crystals were vacuum-dried and kept in the dark and cold until
they were ready for use.
Results
Stoichiometry. The reaction involved successive oxidation
of the sulfur center to sulfate and a urea residue:
2-
2ClO2- + PhN(H)CSNH2 + H2O f SO4
+
PhN(H)CONH2 + 2Cl- + 2H+ (R1)
Stoichiometry R1 was achieved in conditions of excess chlorite;
with an oxidant-to-reductant ratio, R ) [ClO2-]0/[PTU]0 g 2.0
and after prolonged incubation of the reaction solutions. The
stoichiometry was also evaluated spectrophotometrically by
observing the absorption peak of PTU at 249 nm and the
chlorine dioxide peak at 360 nm. Figure 1b shows a series of
spectra taken at various ratios of chlorite to PTU. Successive
increases in chlorite concentrations for fixed [PTU]0 showed a
gradual decrease in this peak until it attained some baseline
absorbance from the product, phenylurea. Spectrum e in Figure
1b is the same as that of phenylurea (spectrum b) in Figure 1a.
Chlorite concentrations in excess of that required by stoichi-
ometry R1 gave chlorine dioxide as a product, and this could
be detected by its absorption peak at 360 nm. Qualitatively, it
could be detected by the deep blue-black coloration it gave with
starch prepared with mercuric iodide. Contrary to other oxy-
halogen oxidations,23 the chlorine dioxide formed could not be
utilized for stoichiometric determinations because the excess
chlorite slowly decomposed in acidic media over several days
to give chlorine dioxide which could not be related to the initial
amount of PTU:24
Figure 2. (A) General absorbance traces at 360 nm of the reaction in
excess chlorite ([ClO2-]0/[PTU]0 > 4.00). The reaction starts with a
short induction period followed by rapid formation of chlorine dioxide.
[PTU]0 ) 1.25 × 10-3 M, [H+]0 ) 0.125 M, [ClO2-]0 ) (a) 5.0 ×
10-3M, (b) 6.25 × 10-3 M, (c) 7.5 × 10-3 M, (d) 8.75 × 10-3 M, and
(e) 1.0 × 10-2 M. (B) Effect of chlorite on the induction period. An
inverse relationship also exists for a wide range of chlorite concentra-
tions. [PTU]0 ) 1.25 × 10-3 M, [H+]0 ) 0.125 M. Note that the
intercept on the chlorite concentration axis (2.38 × 10-3 M), which
represents chlorite concentrations needed to consume PTU, confirms
stoichiometry R1 within limits of experimental error (theoretically, 2.5
× 10-3 M needed to consume 1.25 × 10-3 M).
5ClO2- + 4H+ f 4ClO2(aq) + 2H2O + Cl- (R2)
Iodometric techniques, however, were utilized in order to
evaluate the excess oxidizing power in conditions of excess
chlorite. The analyses gave a stoichiometric ratio slightly greater
than the 2:1 given in reaction R1, especially in low acid
conditions. This could be attributed to the reaction of chlorite
with HOCl to give inert chlorate25,26
:
as by iodometric and titrimetric techniques. Stoichiometry R4,
however, was obtained after prolonged incubation periods of
up to 48 h.
ClO2- + HOCl f ClO3- + Cl- + H+
(R3)
Pathway R3 is insignificant in highly acidic environments where
reaction R2 is dominant.27 Although reaction R3 deactivated
some oxychlorine species, reaction R2, however, does not alter
the overall oxidizing power of the product solution, and
iodometric techniques could still be able to evaluate the excess
ClO2-. The stoichiometry of the chlorine dioxide-PTU reaction
was experimentally determined to be
Reaction Kinetics. The reaction dynamics were extremely
complex, with a highly complex dependence on the pH of the
reaction environment. The reaction dynamics were less complex
in conditions of high excess chlorite, in which the oxidant-to-
reductant ratio, R g 4, and in high acid concentrations, [H+]0
g 0.10 M. Under these conditions the reaction showed a clean
induction period before formation of chlorine dioxide. Figure
2a shows such traces at 360 nm. The rate of formation of
chlorine dioxide after the induction period is directly propor-
tional to the chlorite concentrations. There is an inverse
dependence of the induction period to the initial chlorite
concentrations (see Figure 2b). This appears to indicate that
8ClO2 + 5PhN(H)CSNH2 + 9H2O f
5SO42- + 5PhN(H)CONH2 + 8Cl- + 18H+ (R4)
This stoichiometry was derived spectrophotometrically as well

1084 J. Phys. Chem. A, Vol. 109, No. 6, 2005
Chigwada et al.
chlorite is involved in the reaction kinetics that consumes PTU
to the first power. The plot in Figure 2b was also utilized to
evaluate reaction stoichiometry R1. By extrapolating to Tind
)
∞ (1/TInd ) 0), one can evaluate the minimum amount of
[ClO2-]0 needed to produce ClO2(aq) as a product. This was
after assuming that with [ClO2-]0 less than that needed for
stoichiometry R1, no chlorine dioxide would be formed. The
data in Figure 2b show that, with [PTU]0 ) 0.00125 M, the
threshold concentration of chlorite needed for the formation of
chlorine dioxide was 0.0024 M, which is very close to the 2:1
ratio obtained in reaction R1.
Acid Dependence. Acid concentrations influenced the induc-
tion period as well as the rate of formation and amount of
chlorine dioxide formed after the induction period. The effect
of acid varied depending on oxidant-to-reductant ratios as well
as the range of acid concentrations used. In highly acidic
conditions, [H+]0 > 0.10 M with 2 e R e 4, acid concentrations
reduced the induction period (see Figure 3a). This effect on the
induction period is not as pronounced as that observed with
chlorite (see Figures 2a,b). These acid concentrations did not
affect the rapid rate of formation of chlorine dioxide at the end
of the induction period. Higher acid concentrations, however,
produced lower concentrations of chlorine dioxide (see, again,
Figure 3a). This was unexpected since it is known that chlorine
dioxide formation is from interactions among oxychlorine
species which are universally accepted as acid-catalyzed. The
depressed concentrations of chlorine dioxide are possible only
if reactions in the reaction mixture that form chlorine dioxide
are retarded by acid or if chlorine dioxide consumption reactions
are acid catalyzed. Surprisingly, also, lower acid concentrations
0.005 M < [H+] < 0.015 M, had no visible effect on the
induction period (see Figures 3b and c). At lower ratios, 2 < R
< 4, there was no sharp and discernible induction period, but
a swift and instant formation of chlorine dioxide. At higher
ratios, R g 10, the reactions showed an invariant induction
period irrespective of changes in acid (Figures 3a and b) and in
PTU concentrations. Figure 3b shows a series of experiments
undertaken at low acid concentrations showing the general
increase in chlorine dioxide formed as acid is increased.
However, between acid concentrations of 0.015 and 0.020 M,
this trend is reversed (Figure 3c) and further increases in acid
concentrations results in production of reduced amounts of
chlorine dioxide. Separation of Figures 3b and c is to enable us
to best display this reversal in trend.
The Chlorite-Phenylthiourea Ratio. There were two ways
of altering oxidant-to-reductant ratio, R. One involves fixing
the initial chlorite concentrations and varying [PTU]0 (see Figure
4), and the other involves altering both [ClO2-]0 and [PTU]0
simultaneously (see Figure 5). In Figure 4, traces c, d, and e
have R < 2; and so, ultimately, at t∞, the final chlorine dioxide
concentrations decayed to zero. The observed chlorine dioxide
formation within the first 30 s of this reaction was only transient.
As expected, higher PTU concentrations gave longer induction
periods and faster rates of formation of chlorine dioxide. This
is to be expected if the reactive species that control the formation
of chlorine dioxide are derived from the oxidation of PTU or
reduction of ClO2-. Figure 5 shows a wider range of ratios; 10
> R > 1.7. Low ratios, e.g., traces a, b, and c in Figure 5, gave
transient formation of chlorine dioxide, while higher ratios gave
a monotonic increase in chlorine dioxide. Figure 5 can be
regarded as an extension of the time scale shown in Figure 4.
Figure 3. (A) Effect of acid concentrations on the reaction profiles
with [ClO2-]0/[PTU]0 ) 4.00. The observed induction period before
formation of chlorine dioxide as well as the rate of formation of chlorine
dioxide become insensitive to acid at high concentrations. [PTU]0 )
1.25 × 10-3M, [ClO2-]0 ) 5.0 × 10-3M, [H+]0 ) (a) 0.1 M, (b) 0.2
M, (c) 0.3 M, (d) 0.4 M, and (e) 0.5 M. (B) The effect of acid
undertaken at high [ClO2-]0/[PTU]0 ratio of 10.0. In this series of
experiments, in low acid concentrations, acid catalyzes formation of
chlorine dioxide, which is the exact opposite trend to the one shown
in Figure 3A at high acid concentrations. The pH of the reaction medium
in this series of experiments is above that of the pKa of chlorous acid.
[PTU]0 ) 2.0 × 10-4 M, [ClO2-]0 ) 2.0 × 10-3 M, [H+]0 ) (a) 0.0025
M, (b) 0.005 M, (c) 0.0075 M, (d) 0.01 M, (e) 0.015 M. (C) This figure
shows the exact acid concentrations for reversal of the acid effect on
chlorine dioxide formation. The figure shows that as pH of reaction
medium goes below 1.7 (0.02 M acid), less and less chlorine dioxide
is formed. [PTU]0 ) 2.0 × 10-4 M, [ClO2-]0 ) 2.0 × 10-3 M, [H+]0
) (a) 0.02 M, (b) 0.025 M, (c) 0.05 M.
Chlorine Dioxide Formation. Previous studies that involved
chlorite oxidations have used the formation of chlorine dioxide
as a means of following the reaction and evaluating kinetics
constants.23 Chlorine dioxide is formed, in oxychlorine-
organosulfur reaction systems, by an extraneous and purely
oxyhalogen reaction that does not involve the substrate that is

S-Oxygenation of Thiocarbamides: I
J. Phys. Chem. A, Vol. 109, No. 6, 2005 1085
intermediates have been completely oxidized. (ii) In large ratios,
R > 4 (which is much higher than the required stoichiometry
of R ) 2), production of ClO2 will overwhelm the small
concentrations of substrate present such that there will be an
immediate formation of ClO2. Transient and complex ClO2
formation can be obtained at low ratios, typically below 4. (iii)
If rate of reaction of chlorine dioxide with the reducing substrate
is slow, then oligooscillatory behavior will be obtained for
reaction solutions with R < 2. A clear illustration of this can
be seen in the data shown in Figure 5. Monotonic ClO2
formation is obtained at ratios above 2.5 for acid concentrations
of 0.125 M. This threshold ratio will change at different pH
conditions. After long incubation times, all experiments with
ratios below 2 will show no final chlorine dioxide accumulations
(see stoichiometry R1 and Figure 5). The observed transient
chlorine dioxide formation at these ratios clearly shows a
delicate balance between reactions that form chlorine dioxide
and those that consume it. The effect of chlorite concentrations
(effectively, this is a variation in ratio, R) on the formation of
chlorine dioxide is shown in the data in Figures 2a and 5. Traces
with R e 2 will eventually have no chlorine dioxide at t∞. For
a fixed amount of acid concentrations, the rate of formation
and transient amount of chlorine dioxide formed is directly
proportional to the chlorite concentrations. This is to be expected
if the formation of chlorine dioxide is mainly through reaction
R5. The effect of acid on the reaction will be complex due to
the fact that acid influences nearly all elementary reactions in
this mixture. In some it is catalytic, and in some, inhibitory.
Different acid effects will be observed at different pH conditions.
The three main features of chlorine dioxide formation all change
with acid: rate of formation of chlorine dioxide, maximum
transient chlorine dioxide formed, and the rate of consumption
of chlorine dioxide after attaining its peak. Ultimately, the
overall global reaction dynamics will not show any specific
dependence on acid. Figure 3a shows the high acid environment
in which acid is inhibitory in the formation of chlorine dioxide.
The ratio used in these data, R ) 4, does not give an
overwhelming excess of the oxidant, and thus the rate of
consumption of ClO2 by PTU and other reductants has to be
factored into the determination of the overall reaction dynamics.
Figure 3b, however, with R ) 10, makes the observed chlorine
dioxide formation a pure oxyhalogen reaction since most of the
reducing organosulfur compound and its intermediates will have
been completely consumed by the time chlorine dioxide is
formed.
Figure 4. Effect of varying PTU]0 on the formation of chlorine dioxide
at fixed chlorite concentrations. Concentration of PTU is directly
proportional to the amount of chlorine dioxide produced and to the
induction period at fixed chlorite concentrations. [ClO2-]0 ) 5.0 ×
10-3 M, [H+]0 ) 0.125 M, [PTU]0 ) (a) 2.5 × 10-4 M, (b) 5.0 × 10-4
M, (c) 7.5 × 10-4 M, (d) 1.0 × 10-3 M, and (e) 1.25 × 10-3 M.
Figure 5. Effect of [ClO2-]0/[PTU]0 ratio on the absorbance traces at
360 nm. At high ratios of 4 and above, there is a monotonic formation
of chlorine dioxide. At lower ratios, there is transient formation of
chlorine dioxide. Trace (b) represents the exact stoichiometric ratio
that does not produce chlorine dioxide at the end of the reaction but
clearly shows transient chlorine dioxide formation. [HClO4]0 ) 0.125
M, [ClO2-]0 ) 2.5 × 10-3 M. [PTU]0 ) (a) 1.5 × 10-3 M, (b) 1.25 ×
10-3 M, (c) 1.0 × 10-3 M, (d) 7.5 × 10-4 M, (e) 5.0 × 10-4 M, (f) 2.5
× 10-4 M.
Oxidations by Chlorine Dioxide. Figures 6a,b, and c show
the reaction of chlorine dioxide with PTU monitored at 360
nm. The reaction displays what appears to be a biphasic
behavior, with, initially, sigmoidal autocatalytic kinetics, fol-
lowed by normal decay (follow trace c in Figure 6a and all
traces in Figures 6b and c). In excess of PTU, only the initial
(autocatalytic) phase is observed (Figure 6a, traces a and b).
Such sigmoidal decay kinetics have been consistently observed
with respect to chlorine dioxide oxidations.30 Figure 6b shows
absorbance traces obtained in excess ClO2. The autocatalytic
phase lasts until a stoichiometric ratio of 4:5 is attained; i.e., 4
mol of ClO2 to 5 mol of PTU. These data are summarized in
Table 1. This ratio indicates that autocatalysis lasts until all the
PTU has been oxidized to the sulfinic acid (see reaction
stoichiometry R21). After this, a slow monotonic decay of
chlorine dioxide is then observed up to the end of the reaction.
A simple experiment was undertaken in which the synthesized
sulfinic acid of PTU, phenylthiourea dioxide, (phenylaminoimi-
nomethanesulfinic acid) was reacted directly with chlorine
being oxidized.27 In nonradical processes, the first reactive
oxyhalogen species formed in chlorite oxidations is hypochlo-
rous acid, HOCl.28,29 HOCl will then react very rapidly with
remaining chlorite to produce chlorine dioxide:27
2ClO2- + HOCl + H+ f 2ClO2(aq, yellow) + Cl- + H2O
(R5)
Three factors will determine whether chlorine dioxide is formed.
(i) if the rate of reaction between the reducing substrate and
chlorine dioxide is fast, then there will be no observation of
chlorine dioxide until the reducing substrate and its oxidation

1086 J. Phys. Chem. A, Vol. 109, No. 6, 2005
Chigwada et al.
TABLE 1: Analysis of Absorbance Data from Figure 6Ba
expected
∆[ClO2]/
[ClO2]t
∆[PTU]
at end of ∆[ClO2]/ for formation deviation
[ClO2]0/M [PTU]0/M autocatalysis ∆[PTU] of sulfinic acid observed
0.0006
0.00025 0.000418
0.728
0.748
0.744
0.748
0.800
0.800
0.800
0.800
7.20%
5.2%
5.6%
5.2%
0.000574 0.00025 0.000387
0.000514 0.00025 0.000328
0.000465 0.00025 0.000278
a This table shows that the reaction slows down after a nearly
quantitative formation of the sulfinic acid, indicating that the reaction
between ClO2(aq) and the sulfinic acid is very slow. This is also the
position when chlorine dioxide formation commences (see Figure 10).
than the reaction of chlorine dioxide with PTU. Previous reports
on chlorine dioxide oxidations had suggested that the initial part
of the reaction involves the formation of ClO2-, with subsequent
oxidations being undertaken by chlorite.20 Addition of chlorite
in ClO2 oxidations should show a catalytic effect. Data in Figure
6c indicate that this, indeed, is the case. Only micromolar
quantities of chlorite were used in this series of experiments,
and yet the effect is quite noticeable. The initial autocatalytic
part of the reaction is the one that is immediately catalyzed
with the rest of the reaction not catalyzed any further. This is
evidence that chlorite is indeed the dominant pathway utilized
by chlorine dioxide oxidations. The parallel nature of the
absorbance traces after addition of chlorite also suggest that
chlorite is not an autocatalytic species since its catalytic effect
is linear and dependent solely on the amount of chlorite anions
added.
Data Collected at 249 nm. Figure 1 shows that PTU has a
sharp and isolated peak in its UV spectrum at 249 nm. The
products of its oxidation, phenylurea, as well as chlorine dioxide,
also absorb at this wavelength, but with lower extinction
coefficients. Data collected at this wavelength can deliver
unambiguous initial rates of consumption of PTU at the
beginning of the reaction before the reaction product starts to
accumulate in the reaction medium. Figure 7a shows a typical
series of runs performed at 249 nm. They show an initial rapid
decrease in absorbance signal which quickly damps out to give
way to a slower decrease in absorbance. This sequence would
appear to be the opposite of that observed in Figures 6a,b, and
c. The damping out of the absorbance decrease is due to the
formation of an intermediate which is long-lived enough to be
able to accumulate. If one compares comparable data from
Figures 7a and 6b, one notices that the onset of damping in
Figure 7a coincides with the end of the autocatalytic phase in
Figure 6b. The very sharp and abrupt halt in absorbance decrease
shown in Figure 7a is due to the final formation of the product
which contributes substantially to the final absorbance observed.
The data in Figure 7a give a linear dependence of the initial
rate of PTU consumption with initial chlorite concentrations
(Figure 7b). Induction period data (see Figure 2b) had also
suggested a first order dependence of the reaction rate to chlorite
concentrations. Direct oxidation of PTU by chlorine dioxide,
monitored at 249 nm, gives traces that display autocatalysis
(Figure 7c). This is in agreement with the data obtained in
Figures 6a-c. Figure 8a shows the same types of absorbance
traces as those shown in Figure 7a but with variable acid
concentrations. Contrary to the data obtained at 360 nm with
chlorine dioxide, the depletion kinetics of PTU are monotonic
and less complex. Figures 8a and 8b show the effects of acid
on the initial rate of consumption of PTU. Low acid concentra-
tions display a saturation at about 0.02 M HClO4 (Figure 8b).
If the acid is further increased, a retardation is observed. Figure
Figure 6. (A) Direct reaction of PTU and chlorine dioxide monitored
at 360 nm with oxidant-to-reductant ratios below stoichiometric
amounts. A single sigmoidal decay for the consumption of chlorine
dioxide is observed for traces (a) and (b) in which the oxidation does
not advance as far as the sulfinic acid. [PTU]0 ) 2.5 × 10-4 M, [H+]0
) 0.125 M, [ClO2]0 ) (a) 2.0 × 10-4 M, (b) 3.0 × 10-4 M, and (c)
4.0 × 10-4 M. (B) Chlorine dioxide oxidation at oxidant-to-reductant
ratios; [ClO2]0/[PTU]0 >2. This shows two distinct steps of an initial
sigmoidal decay followed by a standard decay. [PTU]0 )2.5 × 10-4
M, [H+]0 ) 0.125 M, [ClO2]0 ) (a) 7 × 10-4 M, (b) ) 8 × 10-4 M,
(c) 9 × 10-4 M, and (d) ) 1 × 10-3 M. (C) Observed effect of chlorite
on the oxidation of PTU by chlorine dioxide. At micromolar quantities,
chlorite accelerates the reaction and maintains two-stage kinetics.
[PTU]0 ) 2.5 × 10-4 M, [ClO2]0 ) 6.76 × 10-4 M, [H+]0 ) 0.125 M,
[ClO2-]0 ) (a) 0, (b) 5 × 10-6 M, (c) 1.5 × 10-5 M, (d) 3.0 × 10-5
M, (e) 5.0 × 10-5 M.
dioxide. This experiment showed that the reaction of the sulfinic
acid and chlorine dioxide was exceedingly slow, much slower

S-Oxygenation of Thiocarbamides: I
J. Phys. Chem. A, Vol. 109, No. 6, 2005 1087
concentrations show a sigmoidal autocatalytic decay of the peak
at 249 nm (not shown in Figure 8c), and a gradual increase in
acid will change the kinetics trace from this convex shape to a
straight line zero order type kinetics (trace a) and finally to the
standard concave decay kinetics curves that are observed in very
high acid environments (traces d-f). The important aspect of
Figure 8c to note is that the second phase of the reaction is
much faster in low acid concentrations, while in high acid
concentrations this intermediate species formed lingers much
longer before forming products.
Chlorine Dioxide Oxidations Monitored at 249 nm. Figures
9a,b show the depletion kinetics of PTU with chlorine dioxide
as the oxidant. Figure 9a shows the dominance of sigmoidal
kinetics even in excess oxidant conditions, although the observed
autocatalysis decreased with increase in oxidant concentrations.
A high excess of oxidant, however, still did not afford pseudo-
first order kinetics. In highly excessive oxidant concentrations,
the depletion kinetics of PTU approached zero order. Figure
9b shows the effect of acid at low acid concentrations. Acid, at
these conditions, 1.3 < pH < 2.7, acid is very effective in
inhibiting the reaction with respect to rate of consumption of
PTU. If acid is increased further, past the concentrations shown
in Figure 9b, its inhibition effect saturates.
Combining Data at 360 nm and at 249 nm. Figure 10
superimposes data collected at 249 nm and at 360 nm. This
shows that chlorine dioxide formation commences at the start
of the second phase of the reaction when monitored at 249 nm.
This is an important experimental observation: whatever
intermediate is formed or accumulated during the second phase
of the reaction; it must be inert to (or reacts very slowly with)
chlorine dioxide (hence its accumulation). This fact, combined
with the experiment that showed a very slow reaction between
chlorine dioxide and phenylthiourea sulfinic acid as well as data
shown in Figure 6b (summarized in Table 1), indicate that this
intermediate species is the sulfinic acid.
Mechanisms
Chlorite-PTU Reaction. We are assuming a nonradical
pathway in which the initial stages of the reaction involve the
formation of the reactive oxyhalogen species, HOCl:
ClO2- + PhN(H)CSNH2 f PhN)CS(OH)NH2 + OCl-; k1
(R6)
followed by
OCl- + H+ ≡ HOCl; Ka2
(R7)
PhNdCS(OH)NH2 represents the unstable sulfenic acid which
can either be oxidized further to a sulfinic acid, PhNdCS(O2H)-
NH2 (R8),31 or can combine with another PTU molecule to form
a stable dimer (R9).
Figure 7. (A) Depletion of PTU monitored at 249 nm for the oxidation
by chlorite. These data also display two-stage kinetics. The products
of PTU oxidation, phenylurea, and chlorine dioxide give the residual
absorbance that can be seen in the absorbance-time traces. [PTU]0 )
1.25 × 10-4 M, [HClO4]0 ) 0.125 M, [ClO2-] ) (a) 4.0 × 10-4 M,
(b) 5.0 × 10-4 M, (c) 6.0 × 10-4 M, (d) 7.0 × 10-4 M, (e) 8.0 × 10-4
M, (f) 9.0 × 10-4 M, and (g) 1.0 × 10-3 M. (B) Relationship between
initial rate of consumption of PTU and initial chlorite concentrations
from the data shown in Figure 7A. There is a solid linear dependence.
[PTU]0 ) 1.25 × 10-4 M, [HClO4]0 ) 0.125 M. (C) Depletion of PTU
in its oxidation by chlorine dioxide monitored at 249 nm. In high excess
chlorine dioxide environments (ratios > 3), a single sigmoidal decay
is observed for the consumption of PTU. This is in agreement with
Figure 6A. [ClO2]0 ) 4.27 × 10-4 M, [H+]0 ) 0.125 M, [PTU]0 ) (a)
2.5 × 10-5 M, (b) 5.0 × 10-5 M, (c) 7.5 × 10-5 M, (d) 1.0 × 10-4 M,
and (e) 1.25 × 10-4 M.
ClO2- + PhN)CS(OH)NH2 f
PhN)CS(O2H)NH2 + OCl- (R8)
PhN)CS(OH)NH2 + PhN(H)CSNH2 f
(PhN)C(NH2)-S)2 + H2O (R9)
The sulfenic acid can also disproportionate into the more
stable PTU and the sulfinic acid:
2PhN)CS(OH)NH2 h
8c summarizes the acid effect by combining a series of acid
concentrations that range from 0.025 to 0.50 M. Very low acid
PhN)CS(O2H)NH2 + PhN(H)CSNH2 (R10)

1088 J. Phys. Chem. A, Vol. 109, No. 6, 2005
Chigwada et al.
Figure 9. (A) Chlorine dioxide oxidation: In stoichiometric excess
chlorine dioxide, a single sigmoidal decay curve for the consumption
of PTU at 249 nm is observed. [PTU]0 ) 1.25 × 10-4 M, [HClO4]0 )
0.125 M. [ClO2]0 ) (a) 1.69 × 10-4 M, (b) 3.38 × 10-4 M, (c) 4.23
× 10-4 M, (d) 5.07 × 10-4 M, (e) 6.76 × 10-4 M, (f) 8.45 × 10-4 M.
(B) Low acid variation in the oxidation of PTU by chlorine dioxide.
In this case, low acid concentrations are inhibitory. All traces show
sigmoidal decay kinetics. [PTU]0 ) 1.25 × 10-4 M; [ClO2]0 ) (a)
3.38 × 10-4 M; [H+]0 ) (a) 0.0025 M, (b) 0.005 M, (c) 0.01 M, (d)
0.015 M, (e) 0.020 M; (f) 0.025 M, (g) 0.05 M.
Literature has several reports of formation of thiosulfinates from
sulfenic acids in the absence of further oxidant.32 After the
initiation reaction R6, the sulfenic acid can only react through
reactions R8-R10. There seems to be no other fate for the
sulfenic acid. In excess oxidant, the favored pathway is R8. In
limited amounts of oxidant (e.g., 1:1 mixture ratios of PTU with
an equivalent of a two-electron oxidant), quantitative dimer
formation can be obtained as in reaction R9. A bulky dimer, of
the type formed in reaction R9, however, due to its expected
low solubility should precipitate from solution. No precipitate
was observed in our experimental observations, thus precluding
reaction R9 as a major route in this mechanism. We expect,
then, that upon depletion of oxidant, reactions of R10-type will
dominate such that the favored stable forms of the sulfur
compound are formed: PTU itself, the sulfinic acid, or sulfate.
In our proposed mechanism, the rate-determining step in the
oxidation of PTU by chlorite is reaction R6. Reaction R6, as
well as any other reactions that involve the reduction of a
chlorine center coupled to the oxidation of a sulfur center, will
be considered irreversible in this mechanism. Reaction R7 is a
rapid protolytic step such that R6 and R7 can be combined as
one reaction even though we recognize that R11 is not a
Figure 8. (A) Absorbance traces at 249 nm showing the effect of acid
on the rate of consumption of PTU at low acid concentrations. Acid
catalyzes the first stage of the reaction but appears to retard the second
stage. [PTU]0 ) 1.25 × 10-4 M, [ClO2-]0 ) 5.0 × 10-4 M, [HClO4]0
) (a) 0.005 M, (b) 0.010 M, (c) 0.015 M, (d) 0.020 M, (e) 0.025 M.
(B) Initial rate plots for low acid concentrations. A maximum rate of
consumption is attained at approximately pH 2.00. Higher acid
concentrations show a retardation of the rate of consumption of PTU.
(c) A combination of high and low acid concentrations and their effects
on the ClO2--PTU reaction. High acid concentrations involve oxidation
by HClO2 and give a fast first step and slower second step (trace f),
while low acid concentrations involve oxidation by both HClO2 and
ClO2-, followed by a much faster second step (trace a). [PTU]0 ) 1.25
× 10-4 M, [ClO2-]0 ) 7.0 × 10-4 M, [HClO4]0 ) (a) 0.025M, (b)
0.05 M, (c) 0.1 M, (d) 0.15 M, (e) 0.20 M, (f) 0.50 M.

S-Oxygenation of Thiocarbamides: I
J. Phys. Chem. A, Vol. 109, No. 6, 2005 1089
The derived rate of reaction with R6 as the single and rate-
determining step will be given by eq 4:
[PTU][Cl(III)]T[H+]
d[PTU]
Rate ) -
) k1
(4)
(1 + Ka1-1[H+])
dt
Our experimental data do not support eq 4 for all the pH ranges
studied. Equation 4 suggests retardation when [H+] . Ka1, as
the term with acid concentration in the denominator will
dominate. The equation also suggests that the inhibitory effect
of acid will diminish at low acid concentrations, pH > pKa1,
but does not support acid catalysis in any pH range. The
experimentally observed acid catalysis effect may have its
genesis from the oxychlorine reactions that produce the reactive
species. Our data would suggest, then, that reaction R12, in
which the protonated chlorous acid reacts directly with PTU,
contributes significantly to the overall reaction rate observed,
and is, in fact, the dominant pathway. Ka1 has been reported33
to be approximately 0.0191 M-1. One would thus expect
saturation in acid effect with no further retardation in conditions
of pH lower than 2.0, which represents the approximate pKa
range of chlorous acid. The rate of reaction can now be derived
using both reactions R6 and R12 to obtain the following rate
law:
Figure 10. Superimposition of traces at 249 and 360 nm. The trace at
360 nm (trace b) had to be amplified 40x to be on the same scale with
the trace at 249 nm. Chlorine dioxide formation commences exactly at
the end of the first stage of the reaction. [PTU]0 ) 1.25 × 10-4 M;
[ClO2-]0 ) 7.0 × 10-4 M; [H+]0 ) 0.050 M.
termolecular process:
ClO2- + PhN(H)CSNH2 + H+ f
PhN)CS(OH)NH2 + HOCl (R11)
[Cl(III)]T[PTU]
d[PTU]
Rate ) -
)
0[k1Ka1 + k2[H+]] (5)
Ka1 + [H+]
dt
R11 is an initial electrophilic attack on PTU followed by
protonation and elimination of HOCl. Our acid dependence data
can be explained by the effectiveness of the direct reaction of
chlorous acid with PTU:
Experimental EValuation of Kinetics Parameters k1 and k2.
The important kinetics parameters in this reaction system are
the rate constants for the reaction between PTU and ClO2- (k1),
HClO2 (k2), and with ClO2(aq). The reaction’s acid dependence,
on its own, was too complex to be utilized for the evaluation
of k1 and k2. Two complementary experimental techniques had
to be used for evaluating k1and k2. The values obtained in these
two methods could then be checked against each other. Equation
5 was utilized for this evaluation by running a series of
experiments at fixed acid concentrations and collecting initial
rates of consumption of PTU at 249 nm. By running a series of
experiments at varying combinations of [PTU]0 and [ClO2-]0,
a plot could be made of initial rate of reaction vs the product
of [PTU]0 and [Cl(III)]T. This plot was preferred over others
(e.g., a [PTU]0 dependence plot or [ClO2-]0 dependence plot)
because it gave the origin as a firm data point. In our statistical
and graphical analysis, we could weight the origin of such a
plot and thus greatly reduce experimentally generated errors.
The slope of the linear part of this plot gave the value of the
term in the brackets of eq 5; but not individual values of k1 and
k2. Next, using as a rough estimate Figure 8c, trace c (for
example) was utilized to evaluate k2 in the limit of high acid
where only the second term in eq 5 was relevant. This value
was brought back into eq 5 as the initial value to then fit the
initial rate data and evaluate unambiguous values of k1 and k2,
with their relevant error ranges. The second complementary set
of experiments involved the use of low acid concentrations in
which both HClO2 and ClO2- species were relevant. Three acid
concentrations were used: 0.001, 0.003, and 0.005 M. For each
of these acid concentrations, six experiments were performed
at specific chlorite concentrations and initial rates were derived
and recorded. Six plots of initial rate vs acid were generated
and k1 (intercept) and k2 (slope) values were evaluated. These
values were then refined by the use of the more accurate data
set from the first series of experiments. The final values adopted
were from this (first) set. This very extensive analysis gave k1
HClO2 + PhN(H)CSNH2 f
PhN)CS(OH)NH2 + HOCl; k2 (R12)
The complex acid dependence data obtained can be explained
on the basis of protonation equilibria for both chlorous acid
and PTU:
ClO2- + H+ h HClO2; Ka1
(R13)
-1
PhN(H)CSNH2 + H+ h [PhN(H)CS(H)NH2]+; Kb (R14)
In reaction R13, Ka1 is the acid dissociation constant for chlorous
acid. Our kinetics data have suggested that, initially, the rate of
depletion of PTU is first order in both PTU and chlorite with a
complex order with respect to acid. If it is assumed that HClO2
is relatively inert (R12 route negligible), then the rate of reaction
can be derived in terms of chlorite and PTU. The total Cl(III)
-
species in solution are partitioned between ClO2 and HClO2
according to the following mass balance equation:
-
[Cl(III)]T ) [ClO2 ] + [HClO2]
(1)
From the following relationship:
[ClO2][H+]
[HClO2] )
(2)
Ka1
chlorite concentrations can be written in terms of the total
chlorine(III) species:
Ka1[Cl(III)]T
Ka1 + [H+]
-
[ClO2 ] )
(3)

1090 J. Phys. Chem. A, Vol. 109, No. 6, 2005
Chigwada et al.
) 31.5 ( 2.3 M-1 s-1 and k2 ) 114 ( 7.0 M-1 s-1. The fact
that k2 > k1 was to be expected since an electrophilic attack by
chlorite on the nucleophilic sulfur atom of PTU will not be as
facile as that by HClO2. The fact that acid was catalytic at low
acid concentrations establishes this fact.
HOCl (reaction R5) or by Cl2 (reaction R16) should be
kinetically indistinguishable.37
2ClO2- + Cl2 f 2ClO2(aq) + 2Cl-
(R16)
The mechanism of reaction R16 was studied by Nicoson and
Margerum38 and proceeds through some intermediate that
involves an adduct of ClO2‚Cl2-, which can then eliminate a
Cl- ion after the transfer of a Cl+ atom to ClO2-. The
mechanism of the oxidation of chlorite by HOCl is also well
known. Taube and Dodgen,37 using isotopic labeling, have
concluded that this oxidation proceeds through the asymmetric
intermediate Cl2O2:26,35,39
Equation 5 supports a saturation of the initial rate as acid is
further increased past the pKa of chlorous acid. Our experimental
observations show a retardation of the rate as acid concentrations
are increased past 0.10 M. To accommodate this retardation,
eq 5 has to be modified for these high acid environments. Under
these conditions, the Cl(III) species will exist predominantly
as HClO2 with negligible amounts of unprotonated chlorite
available. Any variations in rate can no longer be derived solely
from the oxychlorine species. The protonation of thiocarbonyl
compounds, like PTU, is well-known (reaction R14).34 The
nucleophilic sulfur center is easily protonated at low pH
conditions. One would then expect substantial protonation of
PTU at the conditions of [H+] > 0.10 M. We can assume that
the protonated PTU species is inert and combine pathways R6,
R12, and the retarding effect of protonation of PTU, R14, to
give an overall rate of reaction:
ClO2- + HOCl + H+ h Cl2O2 + H2O
Further reaction of the intermediate will yield ClO2:
ClO2- + Cl2O2 f 2ClO2 + Cl-
(R17)
(R18)
Both R17 and R18 are not elementary single steps although
the kinetics of these reactions can be derived from their
molecularities.
It does appear that the rate of formation of HOCl from the
reduction of ClO2- will determine the rate of formation of ClO2.
Composite reaction R11 (R6 + R7) is the only source of HOCl
at the beginning of the reaction. As the reaction proceeds, the
reaction of chlorite with sulfenic acid (produced in reaction R11)
and other sulfur oxo-acids that are not oxidatively saturated will
also yield HOCl.
d[PTU]
Rate ) -
)
dt
[Cl(III)]T[PTU]0
(1 + Kb[H+])(Ka1 + [H+])
[k1Ka1 + k2[H+]] (6)
Equation 6 assumes negligible reaction between chlorous acid
and protonated phenylthiourea. A more complete rate law would
include the reaction between chlorite anion and the protonated
Consumption of Chlorine Dioxide. Figure 5 shows that there
is a delicate balance between the reactions that form chlorine
dioxide and those that consume it. After eliminating R16, we
can approximate that the sole reaction responsible for the
formation of chlorine dioxide is reaction R5. Chlorine dioxide
can be consumed by any of the reducing sulfur species in
solution, but, due to its radical nature (unpaired electron), the
most significant route is its direct reaction with PTU. Our
kinetics data suggest first-order dependence in PTU and chlorine
dioxide. The sigmoidal decay kinetics of chlorine dioxide
oxidations (Figures 6a-c) suggest that the initial stages of the
reaction involve the formation of the reactive species which
then carries the bulk of the oxidation. This appears, in highly
acidic conditions, to be HClO2, which will ultimately produce
HOCl. The radical chlorine center in chlorine dioxide is electron
deficient with a strong formal positive charge, while the sulfur
center in PTU is electron rich. An equilibrium is set up between
chlorine dioxide and PTU with the formation of a loose adduct
held together by a weak coordinate bond between the chlorine
and sulfur centers:20,40,41
phenylthiourea, but this pathway would be negligible in this
-
scheme because high pH conditions would increase [ClO2
]
while decreasing the protonated thiol concentration, while low
pH conditions would decrease [ClO2-] while increasing the
protonated thiocarbamide (PTU). Equation 6 would deliver the
observed complex acid dependence behavior of initial acid
catalysis followed by saturation and then retardation. In highly
acidic conditions, eq 6 can be utilized with very little loss in
precision, without the first term in the square bracket and the
second bracketed term in the denominator which involve
reaction by chlorite. This simplification of eq 6 was used in
the evaluation of k2 for data shown in Figure 8c where it was
assumed [Cl(III)]T ≈ [HClO2].
The observed variations of rate with acid concentrations
depend on the relative magnitudes of k1 and k2. Our experimental
observations indicate acid catalysis up to approximately pH 2.00
followed by a saturation and retardation as acid concentrations
are increased past this point. This would indicate a dominance
of k2 with the further retardation upon saturation of [HClO2]
being effected by the 1 + Kb[H+] term, which represents
protonation of the thiocarbonyl group. Equations 5 and 6 predict
this trend.
ClO2 + S)C(NH2)N(H)Ph h [O2Cl--S)C(NH2)N(H)Ph]
(R19)
Formation of Chlorine Dioxide. Chlorine dioxide is formed
by the one-electron oxidation of chlorite. Two possible oxidizing
species in this reaction medium are hypochlorous acid (see
reaction R5)26,27 and molecular chlorine.35 The standard dis-
proportionation reactions that can produce chlorine dioxide are
too slow to be effective on this time scale. The kinetics of the
hydrolysis of chlorine have been studied by temperature-jump
spectrophotometry, and the kinetics parameters are well known:36
This adduct can either return to the original reactants or react
further with another chlorine dioxide molecule to form chlorous
acid and the sulfenic acid after a hydrolysis:
[O2Cl--S)C(NH2)N(H)Ph] + ClO2(aq) f
PhN)CS(OH)NH2 + 2HClO2 (R20)
The rate-determining step is the formation of the adduct with
R20 being relatively facile, irreversible, and entropy-driven. The
irreversibility of R20 is supported by the reaction’s inhibition
by acid (see Figure 9b). Acid inhibition also proves that R19 is
Cl2 + H2O h HOCl + H+ + Cl-
(R15)
This is a rapid equilibrium, and so the oxidation of chlorite by

S-Oxygenation of Thiocarbamides: I
J. Phys. Chem. A, Vol. 109, No. 6, 2005 1091
the rate-determining step since protonation of PTU would inhibit
formation of the adduct by forming a poorer nucleophile, [PhN-
(H)CS(H)NH2]+. If the protonated PTU is inert, then the rate
of oxidation of PTU by chlorine dioxide will be given by:
PhN)CS(O2H)NH2 + HOCl f
PhN)CS(O3H)NH2 + H+ + Cl- (R22)
PhN)CS(O3H)NH2 + HOCl + H2O f
PhN(H)CONH2 + SO42- + 3H+ + Cl- (R23)
k3[PTU][ClO2]
1 + Kb[H +]
d[PTU]
-
)
(7)
dt
The accumulation of the sulfinic acid as an intermediate
indicates that reactions R22 and R23 must be slower than R6,
R8, and R12. Work by Makarov et al.43 and Otoikhian et al.44
proved that the oxidation of sulfinic and sulfonic acids
predominantly goes through the initial cleavage of the C-S bond
to give sulfoxylate anion and bisulfite respectively:
In the limit of low acid concentrations, k3 was evaluated from
data of the type shown in Figures 6 and 9. The value obtained
of k3 ) 74 ( 22 M-1 s-1 was not very reliable due to
uncertainties brought about by the sigmoidal decay-type kinetics
as well as the uncertainty in the estimation of Kb.
Nonlinear Consumption and Formation of Chlorine Dioxide.
Data in Figures 6a,b, and c display biphasic kinetics with an
initially sigmoidal consumption of chlorine dioxide followed
by a slower monotonic decay. While autocatalysis in chlorine
dioxide formation has been rationalized through the formation
of HOCl, this alone cannot explain the biphasic kinetics. Table
1 shows an analysis of the data in Figure 6b. The changeover
from rapid autocatalytic kinetics occurred after the almost
quantitative formation of the sulfinic acid of PTU:
PhN)CS(O2H)NH2 + H2O f
PhN(H)CONH2 + HSO2- + H+ (R24)
PhN)CS(O3H)NH2 + H2O f
PhN(H)CONH2 + HSO3- + H+ (R25)
Further oxidation of the sulfur leaving groups should be rapid
and nearly diffusion-controlled. These sulfur oxo-acids exist in
the form of resonance-stabilized zwitterions that are much more
stable in acid than in base.43,45 This explains why the sulfinic
acid intermediate is slower in reacting to form products in high
acids but quickly gives products in low acid (see Figure 8c).
The cleavage of the C-S bond in sulfinic and sulfonic acids in
low acid environments is the basis for guanidylation.46
4ClO2(aq) + 5PhN(H)CSNH2 + 2H2O f
5PhN)CS(O2H)NH2 + 4Cl- + 4H+ (R21)
This suggests that chlorine dioxide is inert to phenylthiourea
sulfinic acid or that this reaction is very slow. This was also
proved by the observation that chlorine dioxide reacts very
slowly with phenylthiourea sulfinic acid (refer to the Experi-
mental Section for the preparation of phenylthiourea sulfinic
acid) to the point of almost being inert. The accuracy of Figure
6b is derived from the fact that these experimental observations
are being performed at 360 nm. Since neither PTU nor
phenylurea absorb at 360 nm, our experiment can follow the
consumption of chlorine dioxide almost exactly. The formation
of the phenylthiourea sulfinic acid at the transition between
sigmoidal kinetics and a normal decay is thus not in dispute.
Figure 10 also confirms our assertion: chlorine dioxide is
formed as soon as the sulfinic acid is formed. The observed
rapid formation of ClO2 in Figure 4 can also be explained
similarly: The final observed rate of formation of ClO2 is from
a combination of the reactions that form ClO2 (reaction R5),
and those that consume it (reactions of type R19 + R20).
Formation of the sulfinic acid, however, decreases the rate of
consumption of ClO2(aq), and hence the observed sudden
increase in the formation of ClO2(aq). An examination of Figure
10 confirms this analysis. The formation of the sulfinic acid,
which is (relatively) inert to ClO2 leaves reactions that form
ClO2 dominant over those that consume it, and hence the instant
formation of ClO2 upon commencement of the second phase
of the reaction. Consumption of the sulfinic acid will then
proceed predominantly through oxidation by HClO2 and HOCl
and very insignificantly by ClO2. Chlorine dioxide’s selective
inertness is well known. In a previous submission from our
laboratories, it was noted that the reaction of chlorite and
formaldehyde delivered almost quantitative formation of chlorine
dioxide based on stoichiometry R2 due to the inertness of
chlorine dioxide to formaldehyde and formic acid.42
OVerall Reaction Network. The chlorite-PTU and chlorine
dioxide-PTU reactions can both be comprehensively explained
by a single reaction network of 29 reactions which are shown
in Table 2. Using relevant assumptions, these reactions can be
severely truncated to only a few for high acid concentrations
where chlorite oxidations are deleted in favor of those by
chlorous acid. The direct reaction between chlorine dioxide and
PTU can be derived by utilizing relevant reactions in Table 2
and deleting the rest.
The tabulated mechanism is divided into five distinct sections.
Reactions M1-M3 represent the rapid protolytic reactions that
are responsible for the observed complex acid dependence.
Reaction M3 is relevant at much higher acid concentrations than
M1 and M2. Reactions M4-M11 are reactions that involve
oxidations by ClO2- and HClO2. The whole reaction scheme is
controlled by reactions M4 and M5 which represent the rate-
determining steps. The set of reactions M12-M15 are the
autocatalytic steps which also include formation of chlorine
dioxide. M13 is the only reaction through which chlorine dioxide
is formed throughout the whole mechanism, and in conditions
of excess oxidant, its rate of formation will be dependent upon
rate of formation of HOCl from the reduction of Cl(III) species.
Reactions M16 - M19 represent the dominant oxidation
pathway. The autocatalytic production of ClO2 transfers the bulk
of oxidation to HOCl, even though, at the beginning of the
reaction, before the build-up of HOCl concentrations, the major
oxidation pathway would have been the series of reactions M4-
M11. Reactions M20-M23 involve oxidation by ClO2. From
the basis of our experimental observations, no reactions have
been included that involve oxidation of the sulfinic and sulfonic
acids by ClO2. The set of four reactions, M24-M27 represent
the hydrolysis of the sulfinic and sulfonic acids. While recent
experimental data indicate that these oxo-acids are more stable
in acidic environments,43,45 this mechanism could not handle
stabilization by acid. Further mechanistic studies need to be
Formation of Phenylurea. Further oxidation of the phenyl-
thiourea dioxide (intermediate species) will yield the sulfonic
acid which can further be oxidized to yield phenylurea:

1092 J. Phys. Chem. A, Vol. 109, No. 6, 2005
Chigwada et al.
TABLE 2: Chlorite-Chlorine Dioxide-Phenylthiourea Reaction Networka
number
reaction
kf; kr
M1
M2
M3
ClO2- + H+ h HClO2
OCl- + H+ h HOCl
1 × 109; 1.91 × 107
1 × 109; 3.24 × 101
1 × 109; 1.12 × 106
PhN(H)CSNH2 + H+ h [PhN(H)CS(H)NH2]+
M4
M5
ClO2- + PhN(H)CSNH2 f PhNdCS(OH)NH2 + OCl-
HClO2 + PhN(H)CSNH2 f PhNdCS(OH)NH2 + HOCl
32
115
M6
M7
M8
ClO2- + PhNdCS(OH)NH2 f PhNdCS(O2H)NH2 + OCl-
ClO2- + PhNdCS(O2H)NH2 f PhNdCS(O3H)NH2 + OCl-
ClO2- + PhNdCS(O3H)NH2 f PhN(H)CONH2 + SO42- + Cl- + 2H+
5 × 103
95
15
M9
M10
M11
HClO2 + PhNdCS(OH)NH2 f PhNdCS(O2H)NH2 + HOCl
HClO2 + PhNdCS(O2H)NH2 f PhNdCS(O3H)NH2 + HOCl
HClO2 + PhNdCS(O3H)NH2 f PhN(H)CONH2 + SO42- + Cl- + 3H+
1 × 103
45
12
M12
M13
M14
M15
ClO2- + HOCl + H+ h Cl2O2 + H2O
1.01 × 106; 0.1
1.5 × 103; 5.5 × 10-6
2 × 103
Cl2O2 + ClO2- h 2ClO2 + Cl-
Cl2O2 + PhN(H)CSNH2 + H2O f PhNdCS(OH)NH2 + 2HOCl
Cl2O2 + PhNdCS(OH)NH2 + H2O f PhNdCS(O2H)NH2 + 2HOCl
HOCl + PhN(H)CSNH2 f PhNdCS(OH)NH2 + H+ + Cl-
HOCl + PhNdCS(OH)NH2 f PhNdCS(O2H)NH2 + H+ + Cl-
HOCl + PhNdCS(O2H)NH2 f PhNdCS(O3H)NH2 + H+ + Cl-
HOCl + PhNdCS(O3H)NH2 f PhN(H)CONH2 + SO42- + H+ + Cl-
1 × 103
M16
M17
M18
M19
5 × 103
2 × 103
100
50
M20
M21
M22
M23
ClO2 + SdC(NH2)N(H)Ph h [O2Cl-SdC(NH2)N(H)Ph]
74
[O2Cl-SdC(NH2)N(H)Ph] + ClO2(aq) f PhNdCS(OH)NH2 + 2HClO2
ClO2 + HOS-C(NH2)dNPh h [O2Cl-(HO)S-C(NH2)dNPh
[O2Cl-(HO)S-C(NH2)dNPh f PhNdCS(O2H)NH2 + 2HClO2
1 × 103
50
1 × 103
500
M24
M25
M26
M27
M28
M29
PhNdCS(O2H)NH2 + H2O f PhN(H)CONH2 + HSO2- + H+
PhNdCS(O3H)NH2 + H2O f PhN(H)CONH2 + HSO3- + H+
HSO2- + HOCl f HSO3- + H+ + Cl-
50
5 × 104
1 × 104
4 × 106
5 × 105
HSO3- + HOCl f SO42- + 2H+ + Cl-
2ClO2 + HSO2- + H2O f HSO3- + 2HClO2
2ClO2 + HSO3- f SO42- + H+ + 2HClO2
a Legend: PhNdCS(OH)NH2, phenylthiourea sulfenic acid; PhNdCS(O2H)NH2, phenylthiourea sulfinic acid; PhNdCS(O3H)NH2, phenylthiourea
sulfonic acid. Rate constants: forward and reverse rate constants (column 3) separated by a semicolon. Note: Reactions M8, M11, and M19 all
require a molecule of solvent water for material balance. Units of rate constants derived from the reaction’s molecularity except where water was
involved, and also in M28 and M29 which are composites of two consecutive steps and considered bimolecular.
performed before a rate law for the decomposition of these oxo-
acids can be derived.
Choice and Establishment of Rate Constants. The first three
reactions, being rapid and protolytic could be assigned any rate
constants as long as they were fast enough so as not to be rate-
determining and not violate the pKa’s of HClO2 and HOCl as
well as the pKb pf the thiocarbamide. Since pKa of chlorous
acid26,47 is 1.72; the forward and reverse rate constants utilized
had to satisfy Ka ) 0.0191. Literature reports pKa of HOCl26,33
as 7.49, and so this was also utilized in the same manner as for
HClO2. Kb for the thiocarbamide was difficult to evaluate, but
a value of 8.9 × 102 M-1 was derived for a similar thiocarb-
amide: guanylthiourea.34 However, protonation of PTU was
only relevant in very high acid environments, while most of
our work was performed in the intermediate acid concentration
ranges where Kb was not relevant. In fact, for most of our
Figure 11. Simulations of the data shown in Figure 7C (traces b and
e) using the mechanism in Table 2. Solid line: experimental; dots:
simulation.
simulations, this reaction was shut down with no loss of
accuracy.
Rate constants for M4 and M5 were derived from this study.
M6 and M9 were all made faster than M4 and M5 respectively,
and this was a reasonable assumption after setting M4 and M5
as rate-determining. Reactions M7, M8, M10, and M11 were
guessed for best fit, but could be ignored if M24 and M25 were
made labile enough. M12 and M13 were derived from literature
values.24,27 Since rate of reaction was controlled by rate of
formation of HOCl, M16 and M17 were set such that they were
faster than M4 and M5. Kinetics parameters for reaction M20
were derived from this study, and subsequently, kM22 was made
slightly faster than kM20. Reactions M26 and M27 are known
to be fast and nearly diffusion controlled, and they were
implemented thus. The autocatalytic reactions M14 and M15
were only relevant in low-to-medium acid concentrations. They
were very important in determining the sigmoidal decay kinetics
observed in Figure 7c. This model could adequately simulate
Figures 2a, 3b, 6a, 7c, 9a, and 9b. While it could simulate the
first part of Figure 8c, the uncertainties in the hydrolysis kinetics
of the sulfinic and sulfonic acids made it difficult to simulate
to second phase of that reaction. Figure 11 shows a simulations
fit to the data shown in Figure 7c. These simulations showed
that reactions M28 and M29 are extremely important in

S-Oxygenation of Thiocarbamides: I
J. Phys. Chem. A, Vol. 109, No. 6, 2005 1093
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ClO2 from ClO2.
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Conclusion
(20) Darkwa, J.; Olojo, R.; Chikwana, E.; Simoyi, R. H. J. Phys. Chem.
A 2004, 108, 5576-5587.
The oxidation of phenylthiourea is complicated by the
formation of two stable intermediates: the sulfinic acid and the
sulfonic acid on the pathway towards total desulfurization to
form phenylurea. Its standard use as a goitrogenic must lie in
its ability to abstract the iodine atom needed by the thyroid
gland.13 Its toxicity, however, must be derived from the
deliberate cleavage of the C-S bond to yield highly reactive
sulfur species that may interact with oxygen to yield reactive
oxygen species. This appears to be a well-accepted route for
the establishment of toxicities in thiocarbamides.48,49 The
physiological environment, which is slightly basic, destabilizes
the oxo-acids in favor of the reduced sulfur species after
cleavage of the C-S bond.43,48
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Acknowledgment. We would like to acknowledge Gus
Lawrence for synthesizing phenylthiourea dioxide. We would
also like to thank Fungai Mukome for running ICPMS for us.
This work was supported by Research Grant Numbers CHE
0137435, CHE 0341152, and CHE 0341769 from the National
Science Foundation.
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