Oxyhalogen - Sulfur Chemistry: Kinetics and Mechanism of Oxidation of Amidinothiourea by Acidified Iodate

Article abstract of DOI:10.1021/jp0367068

The oxidation of guanylthiourea, GTU, by mildly acidic iodate and molecular iodine has been studied Its reaction with iodate shows an oligooscillatory formation and consumption of iodine in batch conditions. The major oxidation product is a ring-cyclized product of guanylthiourea, 3,5-diamino-1,2,4-thiadiazole (GTUC), in which the thioureado moiety is oxidized to the unstable sulfenic acid that instantly attacks the distal amino group, eliminating water and forming the five-membered thiadiazole group. In excess GTU conditions, the stoichiometry of the reaction was 1:3 without any further oxidation past a 2-electron transfer, IO₃^- + 3H₂NC(=NH)NH(C=S)NH₂ → I^- + 3GTU-C + 3H₂O, whereas in excess iodate conditions the stoichiometry is 2IO₃^- + 5H₂NC(=NH)NH(C=S)NH₂ + 2H^- → 5GTU-C + I₂(aq) + 6H₂O. In high acid environments the reaction shows two peaks in iodine concentrations in batch conditions, and at lower acid concentrations one observes an induction period followed by a monotonic formation of iodine according to the Dushman reaction. The overall reaction is heavily catalyzed by iodide. The direct reaction of iodine and GTU is fast, with a bimolecular rate constant of (1.10 ± 0.20) × 10⁴ M^-1 s^-1. This reaction is autoinhibitory with the product, iodide, inhibiting the reaction by forming the relatively inert I₃^- species. Acid also inhibits the oxidation of GTU by iodine by protonating the thiol group, thereby reducing its nucleophilicity. A simple mechanistic scheme comprising 9 elementary and composite reactions was found to be adequate in explaining the full reaction scheme.

Full text of DOI:10.1021/jp0367068

1024
J. Phys. Chem. A 2004, 108, 1024-1032
Oxyhalogen-Sulfur Chemistry: Kinetics and Mechanism of Oxidation of Amidinothiourea
by Acidified Iodate1
Edward Chikwana and Reuben H. Simoyi*
Department of Chemistry, Portland State UniVersity, Portland, Oregon 97207-0751
ReceiVed: September 10, 2003; In Final Form: NoVember 2, 2003
The oxidation of guanylthiourea, GTU, by mildly acidic iodate and molecular iodine has been studied. Its
reaction with iodate shows an oligooscillatory formation and consumption of iodine in batch conditions. The
major oxidation product is a ring-cyclized product of guanylthiourea, 3,5-diamino-1,2,4-thiadiazole (GTU-
C), in which the thioureado moiety is oxidized to the unstable sulfenic acid that instantly attacks the distal
amino group, eliminating water and forming the five-membered thiadiazole group. In excess GTU conditions,
-
the stoichiometry of the reaction was 1:3 without any further oxidation past a 2-electron transfer, IO3
+
3H2NC(dNH)NH(CdS)NH2 f I- + 3GTU-C + 3H2O, whereas in excess iodate conditions the stoichiometry
-
is 2IO3 + 5H2NC(dNH)NH(CdS)NH2 + 2H+ f 5GTU-C + I2(aq) + 6H2O. In high acid environments
the reaction shows two peaks in iodine concentrations in batch conditions, and at lower acid concentrations
one observes an induction period followed by a monotonic formation of iodine according to the Dushman
reaction. The overall reaction is heavily catalyzed by iodide. The direct reaction of iodine and GTU is fast,
with a bimolecular rate constant of (1.10 ( 0.20) × 104 M-1 s-1. This reaction is autoinhibitory with the
product, iodide, inhibiting the reaction by forming the relatively inert I3- species. Acid also inhibits the oxidation
of GTU by iodine by protonating the thiol group, thereby reducing its nucleophilicity. A simple mechanistic
scheme comprising 9 elementary and composite reactions was found to be adequate in explaining the full
reaction scheme.
Introduction
oxidative processes, physiologically, are aided by standard
enzymatic systems such as the P-450 groups and the flavin-
Thiols and thiocarbamides, in general, are very important
physiological molecules. Thiourea and its derivatives form a
vast group of highly reactive and physiologically active
compounds.2 In general, the substituted thioureas are toxic.3
Animal studies on the toxicity of thiourea have shown that when
it is administered in drinking water it induces thyroid adenomas
and carcinomas in rats.4 No one knows the origin of tumori-
genicity in this series of compounds, but speculation is that these
physiological effects can be rationalized from their reactivities.5
Most studies on the effect of thioureas on human health have
been reported without any mechanistic rationalization of these
effects.
We recently embarked on a series of experiments aimed at
rationalizing observed physiological effects of several organo-
sulfur compounds on the basis of their reactivities.6 Thiols are
strongly implicated as antioxidants in human health although
the mechanistic basis for such assertions has never been
established.7 Antioxidants are needed to prevent the formation
of and oppose the actions of reactive oxygen species which are
generated in vivo and cause damage to DNA, lipids and proteins.
All indications are that metabolic processes of all organic sulfur
compounds are oxidative.8 This is a reasonable assumption.
Sulfur in most organic compounds is nucleophilic, and these
centers are susceptible to metabolic oxidation, especially S-
oxygenation in which a series of successive sulfur oxo-acids
are formed.9 These sulfur oxo-acids, sulfenic, sulfinic and
sulfonic acids, are physiologically active, and most physiological
effects of the parent compounds can be ascribed to their
formations. For example, the teratogen thioacetamide’s physi-
ological effects can be traced to the formation of the S,S-dioxide
formed during the metabolism of thioacetamide.10 Most of the
containing monooxygenases.11,12 These enzymatic systems,
however, cannot aid a reaction that is not intrinsically feasible.
S-Oxygenation studies have proved to be very useful in
predicting the toxicities of several thiocarbamides. Very similar
thiols and thiocarbamides can physiologically express vastly
varying degrees of toxicities. For example, though phenylthio-
urea is extremely toxic to rats, surprisingly, diphenylthiourea
is innocuous.13-15 The metabolites of phenylthiourea include
inorganic sulfate and an organic urea residue, and those of
diphenylthiourea are ring-cyclized metabolites with an intact
thione group.13 This has led to the general conjecture that,
maybe, the degree of desulfurization can be related to the
attendant toxicity.16 Bolstering this assertion has been a series
of experiments that show that the cleavage of a C-S bond can
initiate the formation of reactive oxygen species from the initial
generation of the sulfoxyl anion-radical which is derived from
the cleavage of this bond.17
Guanylthiourea, GTU, is an extremely important industrial
and biological molecule. In conjuction with mercaptobenzo-
thiazyl disulfide, it is used as an accelerator in sulfur vulcaniza-
tion of natural rubber. In inorganic synthesis its special structure
has been used to synthesize anion-caged supermolecular com-
pounds. It is in the field of medicine that it has been found to
10.1021/jp0367068 CCC: $27.50 © 2004 American Chemical Society
Published on Web 01/17/2004

Oxidation of Amidinothiourea by Acidified Iodate
J. Phys. Chem. A, Vol. 108, No. 6, 2004 1025
be extremely important. It is a well-known stimulator of
intestinal peristalsis and has been experimentally used for the
clinical treatment of bowel paresis in peritonitis.18 It has also
been very promising in clinical trials as an immunostimulant
and tumor cell inhibitor.19,20 Generally, compounds incorporating
the thioureido and guanidino structures, as GTU, are useful as
radioprotective agents.21 GTU and its derivatives gave a 66%
survival rate in lethal radiation doses.
Very few studies had been performed on the oxidation
mechanisms of GTU. A few studies had addressed the reaction
of GTU with hydrogen peroxide but not on a mechanistic basis.
We believe most of these varied metabolic effects expressed
by GTU are derived from its oxidative metabolites. If we can
evaluate its oxidation mechanism, we may also be able to predict
and match substituted GTU’s with their physiological effects.
Our previous studies on oxyhalogen-sulfur chemistry had
clearly indicated that the metabolic route through the sulfur oxo-
acids was dominant.22 We chose the oxidants acidic iodate and
iodine because of their mildness and the fact that the reaction
of the thiocarbamide moiety with molecular iodine has been
implicated as the most important process in the goitrogenic
effects of many thiocarbamides.4
could be qualitatively tested by BaCl2. Iodine formed in excess
IO3 was also determined separately by its absorbance at 460
-
nm (ꢀ ) 770 M-1 cm-1). The stoichiometry of the direct I2-
GTU reaction was determined spectrophotometrically and by
titration. Standardized iodine solution was titrated from a buret
into a GTU solution of known strength. The end point of the
titration could be detected as the point where the iodine color
lingers. This end point could also be enhanced by starch
indicator.
Results
Stoichiometry. The stoichiometry of the reaction was 1:3
with no variations: 1 mol of IO3- to 3 mol of GTU. This would
suggest a 2-electron oxidation of the sulfur center in GTU (the
most nucleophilic part of the molecule). Such an oxidation
would give the well-known unstable sulfur oxo-acid: sulfenic
acid. Without adequate steric factors, we do not expect the
sulfenic acid to be stable.
IO3- + 3H2NC(dNH)NH(CdS)NH2 f
I- + 3H2NC(dNH)NHC(dNH)CSOH (R1)
The sulfenic acid will either dimerize (in the presence of excess
GTU)
Experimental Section
Reagents. Doubly distilled water was used for preparation
of all stock solutions. The following reagents were used without
further purification: guanylthiourea, GTU, 97% (Acros), potas-
sium iodate (Aldrich), potassium iodide, sodium perchlorate,
and perchloric acid (70%) (Fisher). Iodine (Aldrich), was used
as the resublimed 99.8% crystalline solid. A saturated solution
of iodine was prepared by adding an excess amount of iodine
crystals and stirring for a 24 h period and then letting the
resulting solution to settle overnight. A clear, saturated solution
was obtained and standardized against thiosulfate with starch
as the indicator. Iodine concentrations were also determined
spectrophotometrically by measuring the iodine absorbance at
460 nm, which was the experimentally determined isosbestic
point. The absorptivity coefficient was found to be 770 M-1
cm-1. It is difficult to dissolve iodine concentrations past 10-3
M; however, the solubility was greatly enhanced by the use of
potassium iodide. Higher aqueous iodine concentrations could
be prepared by using excess potassium iodide concentrations
or carefully measured ratios of iodine to KI. The stock solutions
of GTU were prepared fresh daily and protected from light by
covering with aluminum foil.
H2NC(dNH)NH(CdS)NH2 +
H2NC(dNH)NHC(dNH)CSOH f
(H2NC(dNH)NHC(dNH)CS-)2 + H2O (R2)
or disproportionate to the more stable sulfinic acid:
2H2NC(dNH)NHC(dNH)CSOH f
H2NC(dNH)NHC(dNH)CSO2H +
H2NC(dNH)NH(CdS)NH2 (R3)
However, because there was no evidence for any further
oxidation past the 2-electron transfer, reaction R3 could not be
justified. Upon running the products through a GC-mass
spectrometer, a number of fragments were observed, with one
corresponding to the dimeric species formed in reaction R2,
above, and a strong predominant peak for a product with a
molecular weight of 116.4. This is the ring-cyclized product;
3,5-diamino-1,2,4-thiadiazole (GTU-C) with the structure:
Methods. Reactions were run at a constant ionic strength of
1.0 M by adding the required amount of sodium perchlorate.
Reaction temperature was maintained at 25 ( 0.1 °C. The two
reaction systems studied, IO3--GTU and I2-GTU reactions,
were found to be too fast for conventional methods; thus all
kinetic studies were performed on a Hi-Tech Scientific Double-
Mixing SF61-DX2 stopped-flow spectrophotometer. Digitization
and amplifying was done via an Omega Engineering DAS-50/1
16-bit A/D board interfaced to a Pentium II 450 MHz computer.
Potassium iodate, perchloric acid, and sodium perchlorate were
premixed in one vessel at double the reactor concentrations as
one of the reagent solutions, with the other solution containing
only GTU.
Stoichiometric Determinations. Stoichiometric determina-
tion of the iodate-GTU reaction was performed both in excess
IO3- and in excess GTU. In the former case the total oxidizing
power of (IO3- + HOI + I2) was determined by titration. Excess
acidified iodide was added to the reaction solution, and the
released iodine was titrated against standard thiosulfate Sulfate
We conclude that the stoichiometry of the reaction in excess
GTU is
IO3- + 3H2NC(dNH)NH(CdS)NH2 f
I- + 3GTU-C + 3H2O (R4)
Excess iodate conditions produced molecular iodine as a product.
The amount of iodine produced was 0.2 times the amount of
GTU used. This was proved spectrophotometrically and was
highly reproducible (see Figure 3a). The oxidizing power of
the reaction solution, using iodometric techniques, however, did
not vary from that obtained from stoichiometry R4. In excess
iodate, this power was distributed into the different iodine
species: iodine, iodate, and the other oxy-iodine species. This
led to the deduction that excess iodate conditions involved the

1026 J. Phys. Chem. A, Vol. 108, No. 6, 2004
Chikwana and Simoyi
Figure 1. Effect of progressively increasing [IO3-] at constant acid
and [GTU]. Oligooscillations in iodine concentrations are observed.
Final iodine concentrations in all traces shown are the same. [GTU]o
) 0.01 M, [H+] ) 0.05 M, [IO3-]o ) (a) 0.01, (b) 0.02, (c) 0.03, (d)
0.04, (e) 0.05 M.
addition of the Dushman reaction23-25 to reaction R4; eliminat-
ing I-:
IO3- + 5I- + 6H+ f 3I2 + 3H2O
(R5)
Addition of 5R4 + R5 gives the observed stoichiometry in
excess iodate conditions with the 5:1 ratio of substrate to iodine
formed:
2IO3- + 5H2NC(dNH)NH(CdS)NH2 + 2H+ f
5GTU-C + I2(aq) + 6H2O (R6)
Figure 2. (a) Effect of [IO3-] variation on its oxidation of GTU for
the initial part of the reaction showing a finite and measurable induction
period before formation of iodine. [GTU]o ) 0.01 M, [H+]o ) 0.02 M,
[IO3-]o ) (a) 0.01, (b) 0.015, (c) 0.02, (d) 0.03, (e) 0.04 M. (b) Linear
plot of reciprocal induction time vs [IO3-]o for data shown in Figure
2a.
Reaction Dynamics. The reaction is very complex and
oligooscillatory in molecular iodine. Figure 1 shows a series of
typical absorbance traces collected at 460 nm to follow the
production of molecular iodine. In these highly acidic conditions
there is an initial rapid formation of iodine followed by its
consumption. All traces shown in Figure 1 were at excess iodate
environments (i.e., ratio [IO3-]0/[GTU]0 > 0.33), and so the
final stoichiometry would be R6 with iodine as one of the
products. Thus, after this transient consumption of iodine, there
is a final iodine formation that will ultimately give stoichiometry
R6. Because the initial substrate concentration, [GTU]0, is fixed,
all traces shown in Figure 1 will finally give the same
concentration of iodine, which in this series of experiments,
was 0.002 M with an absorbance of approximately 1.55.
Surprisingly, higher iodate concentrations supported a much
more rapid rate of transient iodine consumption (see trend in
traces a-e in Figure 1, ca. 100 s). (This would be unexpected
if it is assumed that the Dushman reaction (R5) is wholly
responsible for the formation of iodine.)
Figure 2a shows a closer examination of the initial stages of
the data shown in Figure 1. Remarkably, one sees a series of
highly reproducible traces with an initial well-defined induction
period followed by a rapid formation of iodine. The induction
period, at pH 2 and below is only a few seconds in duration.
Iodate concentrations, as shown in Figure 2a, appear to control
not only the induction period but also the rate of formation of
iodine after the induction period. One can assume that the end
of the induction period represents the point at which the
reactions forming iodine (e.g., (R5)) supersede those that
consume it. In a sense, the induction period can be used to
estimate the rate of reaction because the inverse of the induction
period should be proportional to the rate of reaction. A plot of
the inverse of the induction time versus initial iodate concentra-
tion gives a straight line (Figure 2b). This indicates that the
rate-determining step is dependent on iodate concentrations to
the first power. The plot in Figure 2b can also be used to
evaluate stoichiometry R4. For as long as GTU is in stoichio-
metric excess, no iodine will be formed, and stoichiometry R4
will dominate. These conditions will not produce iodine at low
acid concentrations, and the induction period will approach
infinity with the inverse of this value being zero. As soon as
iodate is in excess, a finite induction time is observed. Figure
2b confirms this stoichiometry by delivering an intercept for
infinite induction time, which gives a limiting stoichiometric
ratio of 1:3 before iodine formation.
Acid merely acts as a very powerful catalyst in the oxidation
of GTU. Figure 3a shows that although acid reduces the
induction period, it does not influence anything else. All traces
shown in this figure, as expected, end up delivering the same
amount of molecular iodine, which is controlled (at excess iodate
concentrations) by the amount of GTU available. With [GTU]0
) 0.01 M, we expect 0.002 M iodine with a final absorbance
of approximately 1.55, as shown in Figure 3a. A plot of
induction period versus the inverse of the square of the acid
concentrations shows a straight line, indicating that the reaction

Oxidation of Amidinothiourea by Acidified Iodate
J. Phys. Chem. A, Vol. 108, No. 6, 2004 1027
Figure 3. (a) Effect of varying acid concentrations for reactions run
in excess iodate. Acid is not a reactant in the reaction but strongly
catalyzes the reaction by reducing the induction period. [GTU]o ) 0.01
M, [IO3-]o ) 0.03 M, [H+]o ) (a) 0.015, (b) 0.016, (c) 0.017, (d) 0.018,
(e) 0.02 M. (b) Plot showing the linear dependence of induction time
on the reciprocal of [H+]2. [GTU]o ) 0.01 M, [IO3-]o ) 0.03 M. The
plot is derived from data shown in Figure 3a.
Figure 4. Absorbance traces for GTU variation in its oxidation by
iodate. Initially, the absorbance increased (a)-(c) and as the stoichio-
metric point (e) is approached they start to decrease (c)-(e). There is
also a direct relationship between the induction period and GTU
-
concentration as expected. [IO3
]
) 0.005 M, [H+]o ) 0.02 M,
o
[GTU]o ) (a) 0.005, (b) 0.0075, (c) 0.01, (d) 0.0125, (e) 0.015 M.
(b) Effect of acid variation at stoichiometric point (trace (e) from
Figure 4a). The effect of acid can be seen in the form of a shorter
induction period and a higher maximum iodine concentration. [GTU]o
) 0.015 M, [IO3-]o ) 0.005 M, [H+]o ) (a) 0.02, (b) 0.022, (c) 0.026,
(d) 0.03 M.
responsible for the formation of iodine is dependent on acid to
the second power (see Figure 3b).
Iodate:GTU Ratios. The most important parameter in
determining reaction dynamics is the oxidant-to-reductant ratio,
R ) [IO3-]0/[GTU]0. At fixed acid concentrations, a variation
of R will show successively, as R is reduced, an induction period
with a monotonic formation of iodine, an induction period with
two peaks in iodine concentrations (see Figure 1), an induction
period with transient formation of iodine (Figure 4a, trace e),
and no iodine formation at all (at very low R values). Traces
a-d in Figure 4a represent excess iodate concentrations with R
> 0.333. Trace e represents the stoichiometric equivalent with
R ) 0.333. Yet, remarkably, it shows transient formation of
iodine even though stoichiometry R4 does not support iodine
formation. Such oligooscillatory behavior26,27 can only be
supported by the existence of nonlinear kinetics, most likely
fueled by autocatalysis in one of the reactions involved in the
reaction network. Taking trace e in Figure 4a and varying acid
at this stoichiometric ratio gives the traces shown in Figure 4b.
High acid concentrations give shorter induction periods, higher
transient iodine concentrations, and higher rates of depletion
of iodine as the reaction heads toward stoichiometry R4.
Eventually, as acid is continuously reduced, no transient
formation of iodine is observed, and the reaction monotonically
attains stoichiometry R4.
Direct Reaction of Iodine and GTU. Apart from the
Dushman reaction,28 the direct oxidation of GTU by iodine is
the most important reaction in this mechanism. The rate of this
oxidation will determine whether GTU and iodine can coexist
-
at the time scales of the GTU-IO3 reaction. If the I2-GTU
reaction is fast, then formation of iodine would denote that all
the GTU has been consumed. Figure 5a shows that the rate of
this reaction is comparable to the iodate-GTU reaction. The
remarkable feature of the iodine-GTU reaction is that it starts
off extremely rapidly but quickly slows itself down about 0.5 s
into the reaction. It would appear that it is autoinhibitory. The
only product of the reduction of iodine is iodide. Addition of
iodide is shown, in Figure 5b, to be inhibitory (as expected).
To enhance the solubility of iodine, each reaction solution shown
in Figure 5b was prepared as an equimolar mixture if I2/I-. For
example, carefully weighed out crystals of iodine were added
to a 0.001 M KI solution for trace a. In this case, the mixture
of 0.001 M each of I2 and I- gave a final species distribution
of [I2] ) [I-] ) 3.36 × 10-4 M and [I3-] ) 6.64 × 10-4 M.
The five traces shown are not very different from each other
due to the fact that, in all cases, there is enough iodide, such

1028 J. Phys. Chem. A, Vol. 108, No. 6, 2004
Chikwana and Simoyi
Figure 5. (a) Varying iodine concentrations for condition of excess iodine. (Iodine solutions were prepared by dissolving iodine crystals in water,
and no iodide was added to enhance solubility.) The reaction is initially very fast but is essentially over within a second in these unbuffered
reaction conditions. [GTU]o ) 0.0001 M, [I2]o ) (a) 2.8 × 10-4 M, (b) 3.4 × 10-4 M, (c) 3.9 × 10-4 M, (d) 4.5 × 10-4 M, (e) 5.0 × 10-4 M.
-
(b) Absorbance traces of the reaction between GTU and I2/I3 showing the effect of progressively increasing iodine concentration. Equimolar
iodine-iodide concentrations were used to enhance solubility of iodine. These series of reactions are considerably slower than those in Figure 5a
without added iodide. The reaction starts nearly as fast as the experiments in Figure 5a but quickly decelerates as more iodide is formed.
[GTU]o ) 0.001 M, [I2]o ) (a) 0.001, (b) 0.00125, (c) 0.0015, (d) 0.00175, (e) 0.002, (f) 0.0025 M. (c) Effect of iodide variation on the oxidation
of GTU by iodine for trace d in Figure 5a. There is a decrease in the initial rate of reaction as iodide concentration is increased and the reaction
takes much longer to go to completion [GTU]o ) 0.0001 M, [I2]o ) 4.5 × 10-4 M; [I-]o ) (a) no iodide, (b) 2.0 × 10-4 M, (c) 3.0 × 10-4 M,
(d) 4.0 × 10-4 M, (e) 5.0 × 10-4 M. (d) Plot of inverse of reaction rate vs iodide concentrations for the data shown in Figure 5c. The plot is
-1
based on eq 3 where the intercept is -Keq and the slope will allow for the evaluation of k13. This plot gives Keq ) 440 ( 40 M with a k13
)
(1.10 ( 0.2) × 104 M-1 s-1
.
that most of the iodine used as oxidant is converted, predomi-
nantly, to triiodide:29
Catalytic Effect of Iodide. Results from the direct reaction
of iodine and GTU show an autoinhibitory effect from iodide
(Figure 5b). Investigations were made into the effect of iodide
ions on the whole reaction (Figure 8). Trace a with no iodide
added (apart from that initially available in iodate solutions)
has a longer induction period before formation of iodine.
Addition of more iodide progressively decreased the induction
period. Addition of iodide also increased the final amount of
iodine obtained (in excess iodate conditions). This is to be
expected if the Dushman reaction28 is a major player in this
mechanism.
I2(aq) + I- h I3 (aq)
Keq ) 770 M-1
(R7)
-
Despite this, a discernible inhibition is still observed with iodide.
-
The inhibition is due to the different reactivities of I2 and I3
.
Figure 6a shows the inhibitory effect of acid on the direct
reaction of iodine with GTU. Within a small range of acid
concentrations, pH 2-3, there is an inverse dependence of the
initial rate of iodine consumption with acid (see Figure 6b).
Figure 7 shows how rapidly the reaction commences before it
autoinhibits. The traces shown are at different GTU concentra-
tions and fixed [I2]0. One expects the initial absorbance observed
at 460 nm to be approximately 1.16 for all traces. The trace
with the lowest GTU concentration gives the expected initial
absorbance due to iodine, but the rest give initial absorbances
that are much less than 1.16. This means, then, that, within the
mixing and “dead time” of our stopped-flow instrument, more
than half of the reaction will have occurred in traces c-e.
Mechanism
The reaction dynamics suggest that the standard oxyiodine
reactions are dominant in this mechanism. The dependence of
the reaction rate on the acid concentrations to the second-order
implicate the well-known first step in the reduction of iodate:24
IO3- + 2H+ + 2e- f HIO2 + OH-
(R8)

Oxidation of Amidinothiourea by Acidified Iodate
J. Phys. Chem. A, Vol. 108, No. 6, 2004 1029
Figure 7. Effect of GTU variation on its oxidation by I2/I3- solutions.
There is faster initial consumption of iodine followed by a rapid shut-
down of the reaction as it proceeds. [I2]o ) 0.0015 M, [GTU]o ) (a)
0.0005, (b) 0.001, (c) 0.0015, (d) 0.002, (e) 0.0025 M.
Figure 6. (a) Effect of varying acid concentration on the reaction of
GTU and iodine. The iodine used here is dissolved in equimolar
solutions with iodide. The initial rate of reaction decreases as the acid
concentration is increased. [GTU]o ) 0.001 M, [I2]o ) 0.001 M, [H+]o
) (a) 0.001, (b) 0.0015, (c) 0.002, (d) 0.0025, (e) 0.003 M. (b) Inverse
acid dependence plot of the initial rate of the I2-GTU reaction for
data shown in Figure 6a. This is the same plot as the one shown in
Figure 8. [I-]o effect on the iodate oxidation of GTU. The effect of
iodide can be seen on the reduced induction period and increased final
absorbance of iodine as iodide concentration is increased. [GTU]o )
0.01 M, [H+]o ) 0.02 M, [IO3-]o ) 0.01 M, [I-]o ) (a) no iodide, (b)
0.00025, (c) 0.0005, (d) 0.00075, (e) 0.001 M.
Figure 5d with the intercept giving Ka-1
.
The two electrons can be supplied by any 2-electron reductant.
This initial step has been extensively studied, and its kinetics
have been evaluated in a composite form as
HIO2 can also disproportionate in the presence of iodide to
produce more HOI:
rate ) k0[IO3 ][S][H+]2
(1)
-
HIO2 + I- + H+ h 2HOI
(R11)
where S is any 2-electron reductant. If S is iodide, then the
rate-limiting step can be written as30
Reactions R9 and R11 have been extensively studied,31 and their
kinetics parameters are known.
Iodine Formation. There is a single reaction that can be
identified in this reaction network as being responsible for the
formation of iodine. It is the reverse of the iodine hydrolysis
reaction:32
IO3- + 2H+ + I- f HIO2 + HOI
(R9)
HIO2 and HOI represent the active oxyiodine species that
perform the bulk of the subsequent oxidations. Iodate itself is
relatively inert. Thus any induction period observed denotes the
time taken for the buildup of the reactive species. Iodate
solutions always contain at least 10-6 M of iodide ions, and
these are normally sufficient to initiate the reaction because only
catalytic amounts are needed. The rest of the iodide ions are
produced as the reaction proceeds. Further reaction of HOI, for
example, immediately produces I-, which can be recycled and
used in reaction R9:
HOI + I- + H+ h I2(aq) + H2O
KH
(R12)
-1
This reaction has to compete with the other reactions that
consume iodine. The predominant route to consumption of
iodine is the direct reaction of iodine with the substrate:
I2 + GTU f GTU-C + 2I- + 2H+
(R13)
Figure 5b shows that this reaction is autoinhibitory, but as long
as there is a sink for iodide, which, in this case, is made up of
HOI + GTU f GTU-C + H2O + H+ + I-
k10 (R10)

1030 J. Phys. Chem. A, Vol. 108, No. 6, 2004
Chikwana and Simoyi
reactions R11 and R12, autoinhibition will not asset itself.
Visible formation of iodine, which denotes the end of the
induction period occurs when reaction R12 exceeds reaction
R13:
an increase in acid would decrease the rate of reaction. In this
case, both I2 and HOI are active, and a plot of inverse acid
concentrations versus initial rate should show a straight line
with a positive intercept (see Figure 6b). The rate of consump-
tion of iodine can then be derived as
d[I2]
) V12 - V13
(2)
-d[I2]
dt
dt
) [GTU](k10[HOI] + k13[I2])
(5)
These reactions are effectively buffered because acid concentra-
tions far exceed the concentrations of the reacting species.
Reaction R12 will depend on the rate of formation of HOI in
reaction R9, hence the indication, in Figure 2b, that the overall
reaction is controlled by acid to the second power.
Iodine Consumption. Figures 5a and 7 suggest that the
reaction of iodine with GTU is first order in both iodine and
GTU. The initial step in this reaction is an electrophilic attack
on the sulfur center of GTU by iodine:
After elimination of HOI, the rate equation can be deduced to
be
-d[I2]
dt
k′
) [GTU][I2] k13
+
(6)
+
-
(
)
[H ][I ]
where k′ ) KH • k10
Using the mass balance equation on the iodine species of
I2 + >CdS h [>CdSsI]+ + I-
This is then followed by a hydrolysis:
(R14)
2[I2]0 ) 2[I2] + [HOI] + [I-]
one can rewrite eq 6 as
(7)
[>CdSsI]+ + H2O f >CsSOH + I- + 2H+ (R15)
k′
k13
+
+
-
{
}
}
-d[I2]
dt
[H ][I ]
KH
>C-SOH represents the transient sulfenic acid that is formed
before the ring cyclization occurs to form 3,5-diamino-1,2,4-
thiadiazole:
) [GTU][I2]0
(8)
1 +
+
-
{
[H ][I ]
>CsSOH f GTU-C + H2O
(R16)
Equation 8 combines oxidation pathways R10 and R13. This
equation can be simulated to anticipate the functional depen-
dence of the rate of reaction with respect to acid, iodide, oxidant,
and reductant. Burger and Liebhafsky33 evaluated KH as 4.13
× 10-12 M2, and substituting this value into eq 8 predicts acid
dependence for pH conditions above 9 and virtually no acid
effects below this pH because at low pH conditions, 1 . KH/
[H+][I-]. This is contrary to our observations. We can conclude
that eq 4 represents the correct rate law in our mechanism and
that at fixed iodide concentrations, reasonable acid retardation
will be observed within the pH range used in this study and the
same applies to the iodide inhibition. Equation 4 is derived from
the assumption that both the triiodide complex and the proto-
nated GTU are inert, or at least react much more sluggishly
than uncomplexed iodine and unprotonated GTU. The equilibria
for the formation of these inert forms are rapid in both directions.
Evaluation of Rate Constants. If we assume constant acid
concentrations, we can rewrite eq 3 in terms of iodide
concentrations as a function of Keq and initial rate of the reaction:
Iodide inhibition would suggest that the triiodide species is inert
toward GTU or at least is a poorer electrophile. If one considers
it inert, then the rate of iodine consumption will be given by
-d[I2] k13[GTU][I2]0
)
(3)
1 + Keq[I-]
dt
where [I2]0 is the initial concentration of iodine. Because Keq,
at 770 M-1, is high, small amounts of iodide can induce a visible
retardation. Equation 3 supports the data shown in Figure 5b,
which predicts a saturation as iodide concentrations are in-
creased.
Acid Dependence. Acid retardation is difficult to rationalize
because there are two reasonable ways in which acid could effect
a retardation of the oxidation of GTU by iodine. The first
possibility is through the deactivation of the sulfur center for
electrophilic attack by acid:
>CdS + H+ f [>CdSsH]+
Ka-1
(R17)
k13 [GTU]0[I2]0
1
Keq
[I-] )
-
(9)
(
)(
)
Keq mod|rate|
This protonation should be rapid in both directions, but it
effectively blocks reaction R14, which represents the electron
transfer from the sulfur center to iodine. Incorporating the
retardation by acid into the overall rate of reaction gives the
following rate equation:
where mod|rate| is the initial rate of the reaction expressed as
-d[I2]/dt. Equation 9 is also valid in unbuffered reaction
solutions at initial conditions. A plot of the initial iodide
concentrations vs the inverse of the initial rate should give a
straight line with intercept Keq-1. This plot is shown as Figure
5d. This plot was first performed without weighting factors in
an effort to see how close to the literature value of Keq these
-d[I2]
dt
k13[GTU]0[I2]0
(1 + Keq[I-])(1 + Ka-1[H+])
)
(4)
series of experiments would deliver. This analysis gave Keq
)
Equation 4 supports the data shown in Figure 6a,b.
440 ( 40 M-1. Later the intercept was weighted at the inverse
of the expected value of 770 M-1 and treatment based on eq 9
gave a value of k13 ) (1.10 ( 0.2) × 104 M-1 s-1. The same
treatment was extended to acid dependence experiments (Figure
6a,b) to evaluate a reasonable value for Ka-1. We could equally
The other kinetically indistinguishable way in which acid
could effect a retardation is through the hydrolysis reaction R12.
High acid would deplete hypoiodous acid, HOI, in favor of
molecular iodine. If HOI is a faster oxidizing agent that I2, then

Oxidation of Amidinothiourea by Acidified Iodate
J. Phys. Chem. A, Vol. 108, No. 6, 2004 1031
TABLE 1: Mechanism of the Iodate-Guanylthiourea
Reaction
rxn
no.
reaction
kf; kr
M1 IO3- + GTU + H+ f GTU-C + HIO2
80
M2 HIO2 + I- + H+ h 2HOI
2.1 × 109, 90
8.6 × 102, 2.0
3.1 × 1012, 2.2
2.8, 1.44 × 103
6.2 × 109, 8.5 × 106
1.1 × 104
M3 IO3- + HOI + H+ h 2HIO2
M4 HOI + I- + H+ h I2(aq) + H2O
M5 IO3- + 2H+ + I- h HIO2 + HOI
M6 I2(aq) + I- h I3-(aq)
M7 I2 + GTU f GTU-C + 2I- + 2H+
M8 HOI + GTU f GTU-C + H2O + H+ + I- 7.5 × 103
M9 GTU + H+ h [GTU-H]+
5.0 × 106, 5.6 × 103
a Legend: GTU, guanylthiourea; GTU-C, 3,5-diamino-1,2,4-thia-
diazole.
rearrange eq 4 after assuming constant iodide (3.36 × 10-4
M
in this set of experiments):
Figure 9. Computer modeling: Traces showing the simulated results
(circles) and the experimental data (solid line) for data shown in Figure
3a. [GTU] ) 1 × 10-2 M; [IO3-] ) 3 × 10-2 M; trace (a) [H+] )
0.015 M; trace (b) [H+] ) 0.02 M.
k13 [GTU]0[I2]0
1
Ka-1
[H+] )
-
(10)
-1
(
)(
)
mod|rate|
Ka
where, in this case, k13 ) k13(1 + Keq[I-])-1, which gave a
numerical value of 0.795k13 for the data shown in Figure 6b.
Using the value of k13 evaluated in eq 9 and data treatment
shown in Figure 5d, from the slope of the graph in Figure 6b
exotic dynamics, as in this reaction, needs nonlinear kinetics in
the form of either autocatalysis, autoinhibition (reactions M6
and M9), or a combination of both as is the case in this reaction.
Computer Simulations. The simple mechanism in Table 1
was simulated using fourth-order Runge-Kutta techniques
derived from the Chemical Kinetics Simulator from IBM’s
Almaden group. Five of the nine reactions (M2-M6)23-25,30,37
have kinetics parameters well-documented in literature, and this
work derived the rate constants for one of the reactions (M7).
Reaction M8 could be estimated from M7 and this left only
two reactions that could be adjusted for best fit: M1 and M9.
-1
we managed to evaluate a value of Ka ) (8.9 ( 1.3) × 102
M-1, which is related to the Kb of the thiocarbonyl group in
GTU. Using only the intercept value of eq 10 gave a slightly
higher value of Ka-1, (9.8 ( 2.2) × 102 M-1. The value of
Ka-1 derived from the slope, however, delivers a more accurate
-1
value for Ka because the intercept amplifies the errors
generated in the slope.
Because we managed to derive some reasonable value for Ka-1
,
Overall Mechanism. The overall reaction mechanism can
be summarized in the form of Table 1. It is a simple reaction
scheme made up of nine reactions distilled from all the possible
reactions. The mechanism is dominated by four well-known oxy-
iodine reactions (M2-M5),24,25,34,35 one rapid I2/I- equilibrium
(M6),29,36 one protolytic reaction (M9), and three irreversible
oxyiodine-GTU reactions (M1, M7, and M8). This simple
mechanism is able to explain and reproduce the complex
oligooscillatory behavior observed in this reaction system. The
rapid reaction M228 can allow us to ignore any reactions that
involved intermediate iodous acid, HIO2, as an oxidant. Reaction
M1 is the initiation step, together with M5. Only trace amounts
of iodide are needed to initiate the reaction because as the
reaction proceeds, there is autocatalytic production of iodide.
This is evident from the composite reaction M1 + M2:
this meant that after guessing a value for kM9, the reverse was
fixed by the equilibrium constant. Reaction M1 became less
important in conditions where small amounts of iodide were
added prior to start of the reaction (Figure 8). Under these
conditions, reaction control was simply based on reaction M5.
Our removal of HIO2 as an active oxidant meant that reaction
M2 became very important and the global reaction dynamics
were very sensitive to the kinetics parameters of M2; even
though literature values are available. For most of our simula-
tions, we used k-M2 of 1.20 × 102 M-1 s-1, despite the fact
that a value of 90 M-1 s-1 is the suggested literature value.
Reaction M6 is known, through laser Raman temperature-jump
experiments29 to be rapid in both directions. However, the use
of these literature values rendered the simulations extremely
stiff, requiring over 12 h on a Pentium IV 2.4 GHz computer
to simulate a typical result shown in Figure 8. Reduction of the
forward and reverse rate constants by a factor of 100 sped up
the simulations without altering the predicted trajectories. Thus,
for most of our simulations, values of kM6 ) 6.2 × 107 M-1
s-1 and k-M6 ) 8.5 × 104 s-1 were used. Despite the simplicity
of the mechanism, it was still able to simulate all the observed
global dynamics of the reaction except some qualitative aspects
of the data shown in Figure 4b. Though the mechanism could
simulate the peak concentrations of iodine and their (period of)
occurrence, it produced peaks that were much sharper than those
that were experimentally observed. However, the mechanism
was able to satisfactorily model data in Figures 2a, 3a, 5a-c,
6a, 7, and 8. Figure 9 shows a simulations fit to the data in
Figure 3a (traces a and e). There is a very reasonable fit to the
experimentally observed induction period, rate of formation of
iodine, the total amount of iodine formed and the acid
dependence. The iodide dependence experiments shown in
IO3- + GTU + 2H+ + I- f GTU-C + 2HOI + H2O (R18)
when one combines reaction R18 with the faster oxidation of
GTU by HOI (reaction M8), then the overall process (M1 +
M2 + M8) shows that one iodide ion will produce 2 iodide
ions as in typical quadratic autocatalysis.
IO3- + 3GTU + I- f 3GTU-C + 2I-+ 3H2O (R19)
The addition of small amounts of iodide (see Figure 8) will
catalyze the reaction by instantly making reaction R9 viable
from t ) 0 without having to build up the concentration of iodide
from stoichiometry R19. However, this catalytic effect quickly
attains a saturation as in excess iodide, reaction R9 will still
remain the rate-determining step with a standard squared
dependence on acid. Figure 8 also shows the diminished catalytic
abilities of iodide as it is steadily increased. The observation of

1032 J. Phys. Chem. A, Vol. 108, No. 6, 2004
Chikwana and Simoyi
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Figure 5c were the easiest to simulate because we could shut
down all the reactions in the mechanism except reactions M6-
M8, with the only tunable parameter being kM7. This also helped
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Dalton Trans. 2000, 511-514.
in establishing our experimental value for kM7
.
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P. M.; Simoyi, R. H. Arch. Biochem. Biophys. 1999, 367, 289-296.
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Acknowledgment. This work was supported by Grant
Number CHE 0137435 from the National Science Foundation.
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