Oxyhalogen-sulfur chemistry: Bromate oxidation of 1-methyl-2-thiourea in acidic medium

Article abstract of DOI:10.1021/jp960259x

The reaction between bromate and 1-methyl-2-thiourea, CH₃NH(NH₂)C=S (MTU), has been studied in acidic medium. The stoichiometry of the reaction in excess MTU has been established as 4BrO₃^- + 3CH₃NH-(NH₂)C=S + 3H₂O → 4Br^- + 3SO₄^2- + 3CH₃NH(NH₂)C=O + 6H⁺ (A); and in excess bromate the stoichiometry is: 8BrO₃^- + 5CH₃NH(NH₂)C=S + H₂O → 4Br₂ + 5SO₄^2- + 5CH₃NH(NH₂)C=O + 2H⁺ (B). Stoichiometry B includes the oxidation of MTU (stoichiometry A) and the oxidation of Br^- by the excess BrO₃^-: BrO₃^- + 5Br^- + 6H⁺ → 3Br₂ + 3H₂O. In excess BrO₃^- the reaction is characterized by an induction period followed by formation of Br₂. The reaction between Br₂ and MTU is very fast with a bimolecular lower limit rate constant of 2.0 × 10⁶ M^-1 s^-1 such that the formation of Br₂ is an indicator for the complete consumption of MTU. The reaction could be followed by monitoring the depletion of MTU through its absorbance at λ = 239 nm or the formation of Br₂ at λ = 390 nm. The rate of reaction was deduced as -d[BrO₃^-]/dt = k₀[BrO₃^-][MTU][H⁺]² with k₀ = 14.7 ± 1.2 M^-3 s^-1. A proposed 17-step reaction mechanism gives good agreement between experimental data and computer simulations.

Full text of DOI:10.1021/jp960259x

J. Phys. Chem. 1996, 100, 13521-13530
13521
Oxyhalogen-Sulfur Chemistry: Bromate Oxidation of 1-Methyl-2-thiourea in Acidic
Medium1
Sreekantha B. Jonnalagadda
Department of Chemistry, UniVersity of Durban-WestVille, PriVate Bag X54001,
Durban 4000, Republic of South Africa
Cordelia R. Chinake and Reuben H. Simoyi*
Center for Nonlinear Science and the Department of Chemistry, West Virginia UniVersity, Box 6045,
Morgantown, West Virginia 26506-6045
ReceiVed: January 24, 1996; In Final Form: May 21, 1996X
The reaction between bromate and 1-methyl-2-thiourea, CH3NH(NH2)CdS (MTU), has been studied in acidic
-
medium. The stoichiometry of the reaction in excess MTU has been established as 4BrO3 + 3CH3NH-
2-
(NH2)CdS + 3H2O f 4Br- + 3SO4 + 3CH3NH(NH2)CdO + 6H+ (A); and in excess bromate the
-
2-
stoichiometry is: 8BrO3 + 5CH3NH(NH2)CdS + H2O f 4Br2 + 5SO4 + 5CH3NH(NH2)CdO + 2H+
(B). Stoichiometry B includes the oxidation of MTU (stoichiometry A) and the oxidation of Br- by the
excess BrO3- : BrO3- + 5Br- + 6H+ f 3Br2 + 3H2O. In excess BrO3- the reaction is characterized by an
induction period followed by formation of Br2. The reaction between Br2 and MTU is very fast with a
bimolecular lower limit rate constant of 2.0 × 106 M-1 s-1 such that the formation of Br2 is an indicator for
the complete consumption of MTU. The reaction could be followed by monitoring the depletion of MTU
through its absorbance at λ ) 239 nm or the formation of Br2 at λ ) 390 nm. The rate of reaction was
deduced as -d[BrO3 ]/dt ) k0[BrO3 ][MTU][H+]2 with k0 ) 14.7 ( 1.2 M-3 s-1. A proposed 17-step reaction
-
-
mechanism gives good agreement between experimental data and computer simulations.
Introduction
kinetics and mechanisms of a sulfur center. The importance
of sulfur chemistry in the environment,7 biology,12 and industry13
made us recently embark on a series of research studies aimed
specifically at kinetics and mechanisms of sulfur oxidations
using, initially, sulfur atoms bonded to small organic mol-
ecules.14
The reactivity of the sulfur center is heavily dependent upon
its environment and its oxidation state. In the free form as S2-
and HS- the sulfur center presents complex and exotic reaction
characteristics especially with respect to its oxidation.2-6 It is
partly because of this complexity that a robust algorithm for
the oxidation of a sulfur atom from an oxidation state of -2 to
+6 has not yet been devised. Efficient environmental abatement
procedures7 cannot be devised until the mechanisms of sulfur
oxidations in both aqueous and gaseous phases have been
deduced.
It is generally accepted that although the sulfur center can
attain several possible oxidation states between -2 to +6, a
special stability is associated with oxidation states -2, 0, and
+6. The 0 oxidation state is especially stabilized by polym-
erization, where S8 is the most stable form of S(0).8 Sulfur
polymerizations will always distort oxidation mechanisms of
the free S(-2) center2 unless it is in the environment of a
powerful oxidizing agent or in overwhelming excess of oxidant.
Sulfur chemistry is greatly simplified if the sulfur center is
bonded to an organic molecule. Some of the reasons for
subsequent simpler kinetics include the following: (a) the C-S
bond is very strong and will not easily release a free sulfur atom,
(b) the bulk of the organic moiety discourages extensive
polymerizations past the dimer stage,9 (c) the cleavage of the
C-S bond can only occur when the sulfur atom is oxidatively
and coordinatively saturated at oxidation state +6,10 (d) oxida-
tion steps are essentially irreversible (in excess oxidant), and
(e) autooxidations and disproportionations, should they occur,
serve only to ensure stoichiometric consistency, especially in
excess of reductant.11
Much insight into sulfur chemistry can be gained from a study
of thiocarbonyl compounds, R1R2CdS. Small molecules such
as thiourea, substituted thioureas, and thiosemicarbazides are
moderately soluble in water, making it possible for the studies
to be carried out in aqueous media.15 The oxidation state of
the sulfur center in thiocarbonyls is -2. Is it possible, then, to
be able to oxidize such a sulfur atom to oxidation state +6
without the inherent complexities associated with sulfur chem-
istry? Would it be possible to cleave the C-S bond, and under
what conditions? What role will the S(0) state play, if any?
Previous studies on the simplest thiocarbonyl, thiourea, had
shown that the CdS double bond is extremely polar,16 much
more so than the corresponding CdO double bond due to the
mismatch of size between the carbon and sulfur atoms which
results in an incomplete π-bond overlap. Thus a permanent
negative dipole resides on the sulfur atom, making it extremely
vulnerable to electrophilic attack.17
We report in this paper on the oxidation of a substituted
thiourea: 1-methyl-2-thiourea, MTU. The oxidation of thiourea
by chlorite18 and by bromate19 exhibits sustained oscillations
in a continuously stirred tank reactor (CSTR). The oxidation
by chlorite further shows lateral instabilities in the form of a
traveling acid wave front.20 A good understanding of the
mechanism of such exotic dynamics requires a good grasp of
the kinetics and mechanism of the “drive” reaction. While
focusing on the sulfur center, it is important to evaluate the
The removal of the major sources of irreproducibilities and
complexities in sulfur chemistry makes it possible to study the
effect of substitution on the rate of oxidation and the ease of
2-
the C-S bond cleavage to form SO4
.
X Abstract published in AdVance ACS Abstracts, July 15, 1996.
Changing the substituents on the nitrogen may affect the rate
S0022-3654(96)00259-6 CCC: $12.00 © 1996 American Chemical Society

13522 J. Phys. Chem., Vol. 100, No. 32, 1996
Jonnalagadda et al.
of reaction but not the mechanism since the major reaction is
at the sulfur center. The study of the kinetics and mechanism
of the oxidation of 1-methyl-2-thiourea should provide support
for the mechanism proposed for the oxidation of thiourea. A
computer simulation of the proposed reaction scheme is also
presented in the paper.
Experimental Section
Materials. 1-Methyl-2-thiourea, MTU, (Aldrich), perchloric
acid, 70-72%, sodium bromate, sodium bromide (Fisher),
bromine, and p-(dimethylamino)benzaldehyde (Aldrich) were
used without further purification. MTU solutions were prepared
just before use and not kept for more than 24 h. Reaction
solutions were prepared using singly distilled water.
Methods. Experiments were carried out at 25 ( 1.0 °C. The
ionic strength was maintained at 1.0 M (NaClO4) in all
experiments with the exception of those run in high acid. The
slower BrO3--MTU reactions were monitored spectrophoto-
metrically by following absorbance of MTU at λ ) 239 nm on
a Perkin-Elmer Lambda 2S UV/vis spectrophotometer interfaced
to a DEC 486-33DX computer. The reaction’s progress could
also be followed by monitoring the absorbance of Br2 at λ )
390 nm. For the faster reactions, kinetics measurements were
performed on a Hi-Tech Scientific SF-61AF stopped-flow
spectrophotometer with an M300 monochromator and a spec-
trascan control unit. The signal from the spectrophotometer
was amplified and digitized via an Omega Engineering DAS-
50/1 16-bit A/D board interfaced to a computer for storage and
analysis. The direct reactions between aqueous bromine and
MTU were also monitored by following decrease in bromine
absorbance at 390 nm on the stopped-flow spectrophotometer.
Stoichiometric Determinations. The stoichiometric deter-
minations were performed both in excess BrO3- concentrations
and excess MTU. In excess BrO3-, the total oxidizing power
(BrO3- + HOBr + Br2) was determined iodometrically. Sulfate
was determined gravimetrically as BaSO4. In these gravimetric
Figure 1. Stoichiometric data: Spectral scans of several reaction
product solutions in which [MTU]0 was steadily increased at constant
[BrO3-]0. Absorbance at 430 nm indicates the amount of complex
formed between the urea product and p-(dimethylamino)benzaldehyde.
The maximum absorbance is obtained at a ratio of 4:3. Any further
increase in [BrO3-]0 did not increase the amount of urea product formed.
product solution with p-(dimethylamino)benzaldehyde. The
absorbance of the resulting solution at 430 nm indicated the
amount of organic product formed. The stoichiometric equiva-
lent was taken as the amount of MTU needed to give maximum
absorbance at 430 nm. Any further increase in [MTU]0 gave
no additional increase in absorbance. Figure 1 shows the
product spectral scans obtained from the reaction mixture, which
show a maximum product formation when ratio R ) [BrO3-]0/
[MTU]0 is 4:3.
-
determinations, BrO3 was first removed iodometrically since
In excess bromate the stoichiometry was determined as
it forms slightly insoluble Ba(BrO3)2 with BaCl2, thus distorting
-
the precipitation results. Bromine formed in excess BrO3
8BrO3- + 5CH3NH(NH2)CdS + H2O f 4Br2 +
5SO42- + 5CH3NH(NH2)CdO + 2H+ (R2)
concentrations could be separately determined by its absorbance
at 390 nm using the absorptivity coefficient of 142 M-1 cm-1
.
Br- was determined quantitatively by precipitation as AgBr
2-
using AgNO3 (after removing SO4 with measured amounts
In overwhelming excess of BrO3-, R > 5, the amount of Br2
formed, [Br2]f, was determined by the [MTU]0 and the ratio of
([MTU]0/[Br2]f) ) 1.25 was consistently maintained (λ ) 390
nm, ꢀ ) 142 M-1 cm-1). One mole of MTU gave one mole of
of BaCl2). The expected organic product, a substituted urea,
was determined spectrophotometrically at λ ) 430 nm by using
p-(dimethylamino)benzaldehyde, which forms a pink com-
plex21 with any compound with an amide group (HNsCdO).
2-
SO4 (BaSO4 precipitation). Working in high acid did not
afford the opportunity to assess the amount of acid formed from
either R1 or R2.
Results
Stoichiometry. The stoichiometry of the reaction in excess
MTU was determined as
The two stoichiometries are essentially the same in terms of
the oxidation of MTU. The only difference lies in the further
oxidation of Br- by BrO3 to produce the bromine in stoichi-
-
4BrO3- + 3CH3NH(NH2)CdS + 3H2O f 4Br- +
3SO42- + 3CH3NH(NH2)CdO + 6H+ (R1)
ometry R2.
Reaction Kinetics. Upon mixing acidic bromate and MTU
solutions, no activity is observed in the leading reaction
indicators: redox potential, absorbance, electrical conductivity,
and pH for a measurable amount of time. With excess MTU
over BrO3-, no activity is observed for an indefinite period,
The stoichiometric ratio of 4:3 was determined in two ways:
(a) A series of experiments were performed at constant MTU
while varying the [BrO3-]0. These reaction solutions were
capped and left overnight. After a 24 h incubation period they
were analyzed for the presence of Br2 (λ ) 390 nm). The
-
although all the BrO3 is consumed and some of the MTU is
oxidized to SO4 and urea according to stoichiometry R1. In
2-
excess BrO3-, however, the reaction starts out with a quiescent
period in which the reaction indicators show no activity. After
some induction period there is a sudden rise in the redox
potential of the solution (Figure 2). Soon after the sudden
increase in redox potential, the solution turns yellowish as
-
stoichiometric equivalent of BrO3 was the amount needed to
just produce Br2 as one of the products. (b) A series of
experiments were also performed at constant [BrO3-]0 with
variable [MTU]0. Excess MTU was determined by reacting the

Oxyhalogen-Sulfur Chemistry
J. Phys. Chem., Vol. 100, No. 32, 1996 13523
Figure 2. Typical redox potential trace for the BrO3--MTU reaction
in excess BrO3-. The rapid increase in redox potential coincides with
the formation of Br2. The final redox potential reading is essentially
the Br-/1/2Br2 couple. [BrO3-]0 ) 0.023 M; [MTU]0 ) 0.0029 M; [H+]0
) 0.143 M.
Figure 3. Spectral scans in the UV region taken every 40 s showing
absorbance activity at 239 nm and no activity anywhere else in the
region. This peak is due to MTU, and its depletion can be monitored
at this wavelength. [BrO3-]0 ) 0.0005 M; [MTU]0 ) 0.000 25 M; [H+]0
) 0.10 M.
aqueous bromine is formed. The redox potential finally
1
stabilizes at a value very close to the potential of the /2Br2/
Br- couple (1.14 V > E > 1.00 V). The length of the induction
period can be adjusted by changing the initial concentrations
of the four major reactant species: acid, MTU, BrO3-, and Br-.
Thus the BrO3--MTU reaction can be regarded as a “clock
reaction”. The induction time, as defined in this paper, is the
time taken by the reaction solution before it produces aqueous
bromine.
formation (λmax ) 390 nm) occurs only after all the MTU has
been consumed. Thus the induction times monitored here
indicate the relative rates of consumption of MTU.
Acid Dependence. The induction times obtained were much
shorter in high acid concentrations (Figure 7a). The data
collected at 390 nm show acid catalysis. Generally, the inverse
of the induction time is a measure of the rate of reaction, and
We have divided the data presentation into four segments:
(a) data collected at 239 nm, (b) data collected at 390 nm before
the induction period, (c) data collected after the induction period,
and (d) data collected on the effect of Br-.
2
in this case a plot of reciprocal induction time vs [H+]0 gives
a straight line (Figure 7b).
Bromate Dependence. Bromate also decreases the induction
period, and a plot of reciprocal induction time vs [BrO3-]0 gives
a straight line (Figure 8).
(a) Data Collected at λ ) 239 nm. The data at this
wavelength were obtained using a conventional UV/vis spec-
trophotometer with a solution blank made up of acidic bromate
in order to subtract the effect of bromate absorption in the UV
range. UV spectral scans for the reaction solution show a peak
at 239 nm, and this can be assigned to MTU. Reaction progress
MTU Dependence. In situations in which 3.0 > R > 1.3,
increases in [MTU]0 will increase the induction period, as it
will take longer to consume MTU and to subsequently trigger
formation of Br2. Figure 9 shows a situation in which R was
maintained at a very high value (300 > R > 50). No change
in the induction period is observed, although there is a change
in the rate of formation of Br2 after the induction period.
(c) Post-Induction Period Data. The majority of these
-
can be monitored in excess BrO3 and in excess MTU by
following the decay of the 239 nm peak. Figure 3 shows a
series of UV spectra for the reaction solution, which show the
decay of the 239 nm peak.
-
Acid Dependence. At constant [BrO3-]0 and [MTU]0, the
reaction is strongly catalyzed by acid (Figure 4a). Stoichiom-
etries R1 and R2 show that acid is not a reactant, but a product.
The catalytic effect of acid must be through the formation of a
experiments were performed in environments such that [BrO3 ]0,
[H+]0 . [MTU]0. By the time the induction period ends,
bromate and acid will still retain concentration values very close
to their initial values. The concentration variables will now be
limited to MTU concentrations and any reactive intermediates
that might be formed. The observation of pseudo-first-order
kinetics in the formation of Br2 after the induction period is as
expected. The data such as those in Figures 7a and 9 were
-
reactive intermediate either with BrO3 or with MTU. A plot
of initial rate of consumption of MTU against [H+]0 gives a
2
straight line with an intercept kinetically indistinguishable from
zero (see Figure 4b).
Bromate Dependence. There is a positive dependence
between initial rate and [BrO3-]0, although the catalytic effect
is not as marked as with [H+]0 (Figure 5a). The initial rate
varies linearly with [BrO3-]0 (Figure 5b) with the expected
intercept through the origin.
analyzed for the apparent pseudo-first-order rate constant, kapp
.
Linear dependence was obtained between kapp and [H+]02 (Figure
-
10a) and [BrO3
] (Figure 10b). In the case of [MTU]
0
dependence, no relationship could be discerned between the
pseudo-first-order rate constant and the initial MTU concentra-
tion. This is expected in a true pseudo-first-order kinetics
environment. The rate constants obtained upon varying [MTU]0
MTU Dependence. Varying initial MTU concentrations
adjusts the initial MTU absorbance (Figure 6a). The reaction
does show a first-order dependence in [MTU] since a plot of
initial rate vs [MTU]0 gives a straight line (Figure 6b).
(b) Data Collected at λ ) 390 nm. Data at λ ) 390 nm
are shown in Table 1. There is a very slight increase in kapp
,
which can be considered insignificant.
(d) Effect of Br-. The effect of bromide was difficult to
-
could only be collected in excess BrO3 conditions. Bromine
assess. It is important to be able to deduce from experimental

13524 J. Phys. Chem., Vol. 100, No. 32, 1996
Jonnalagadda et al.
Figure 5. (a, top) Effect of [BrO3-]0 on the MTU depletion. [MTU]0
) 0.000 25 M. [H+]0 ) 0.30 M. [BrO3-]0 ) (a) 0.015 M; (b) 0.025 M.
(b, bottom) Linear plot of initial rate of MTU consumption vs [BrO3-]0.
The plot also gives the expected zero intercept (reaction will not proceed
without BrO3-). Conditions are the same as in part a.
Figure 4. (a, top) Absorbance traces at λ ) 239 nm showing the effect
of acid. The first 20 s were not recorded, as they can be attributed to
-
sample mixing and preparation. [BrO3
] ) 0.025 M; [MTU]0 )
0
0.000 25 M. [H+]0 ) (a) 0.15 M; (b) 0.20 M; (c) 0.25 M. (b, bottom)
Plot of initial rate of MTU depletion vs [H+]02 from the data in part a.
The intercept, as expected, is slightly negative (some threshold acid
concentration is needed for the reaction to commence).
stopped-flow spectrophotometer. Figure 12 also shows that by
the time this reaction is detected by the photomultiplier most
of the Br2 would have already been consumed. Attempts to
slow the reaction by using Br3- instead of Br2 were unsuccess-
ful. None of the standard dependence experiments could be
performed on this reaction. The important information obtained
from Figure 12 was that the reaction is too fast to be
rate-determining in the BrO3--MTU reaction network. We
managed to deduce a lower limit bimolecular rate constant for
data whether Br- is acting as a general reagent or as a catalyst.
Br- reduced the induction time when [MTU]0 > [Br-]0.
However, when [Br-]0 > [MTU]0 not only did the induction
time decrease but the amount of Br2 formed also increased with
[Br-]0 (Figure 11a). The added Br- ultimately ends up being
converted to Br2. The plot of reciprocal induction time vs [Br-]0
showed a complex plot: below a certain threshold [Br-]0 value,
the induction time does not change. After the threshold value
has been reached, there is a linear dependence. Due to the very
widely spaced [Br-]0 used, a semilog plot was used (Figure
11b) in order to plot all the initial Br- concentrations on the
same plot. Figure 11b shows the catalytic effect of Br- as the
induction period is greatly reduced in high [Br-]0 conditions.
The response of the pseudo-first-order rate constant was also
complex (Figure 11c). The pseudo-first-order rate constant is
unaffected by very low [Br-]0 (e.g. 10-6-10-4 M Br-).
Direct Br2-MTU Reaction. Figure 12 shows that this
reaction is very fast: fast enough to test the limits of our
the Br2-MTU reaction of 2.0 × 106 M-1 s-1
.
Mechanism
The stoichiometric analysis results simplified the search for
a plausible mechanism for this reaction: (a) the organic
molecule is left intact, and the only activity in terms of oxidation
(electron transfer and/or hydrolysis) is occurring at the sulfur
center of the thiocarbonyl compound, MTU; (b) there is
successive electron transfer from the sulfur center to the bromine
center except when the oxidation state of the sulfur center is at

Oxyhalogen-Sulfur Chemistry
J. Phys. Chem., Vol. 100, No. 32, 1996 13525
Figure 6. (a, top) Effect of varying [MTU]0 on the traces obtained at
λ ) 239 nm (absorption peak for MTU). There is an expected decrease
in initial absorbance as the initial MTU concentration is decreased.
[H+]0 ) 0.25 M. [BrO3-]0 ) 0.025 M. [MTU]0 ) (a) 0.000 373 M;
(b) 0.000 25 M; (c) 0.000 188 M; (d) 0.000 125 M. (b, bottom) Plot of
initial reaction rate vs [MTU]0 for the data in part a.
Figure 7. (a, top) Absorbance traces at λ ) 390 nm in conditions of
excess [BrO3-]0. Acid simultaneously reduces the induction time and
accelerates rate of formation of Br2. [BrO3-]0 ) 0.10 M; [MTU]0 )
0.001 M. [H+]0 ) (a) 0.10 M; (b) 0.20 M; (c) 0.30 M; (d) 0.40 M; (e)
0.50 M. (b, bottom) Plot of reciprocal of induction time (equivalent to
2
rate of reaction) vs [H+]0 from the data shown in part a.
+6 and that of the bromine center is at -1; and (c) the
stoichiometry is so clean that it accounts for all the sulfur in
the reaction mixture going to SO42-, thus precluding the
possibility of sulfur polymerizations and/or the formation of
colloidal sulfur.
Reaction R3 does not give the observed second-order depen-
dence on acid concentration. The bromous acid formed in
reaction R3 can further be reduced to form Br-.
HBrO2 + CH3NH(NH2)CdS f
From the data presented in Figures 4b, 5b, and 6b, the rate
of reaction is given by the equation
CH3NH(NH)C-S-OH + HOBr (R4)
-
rate ) k0[BrO3 ][H+]2[MTU]
(1)
HOBr + CH3NH(NH2)CdS f
CH3NH(NH)C-S-OH + Br- + H+ (R5)
From the experimental data the value of k0 was determined to
be 14.7 ( 1.2 M-3 s-1. A rate-determining step that can support
eq 1 is not easily deduced from the expected reaction initiation
and subsequent reaction sequence. The reaction initiation can
be presented through reaction R3:
The Br- produced in reaction R5 can then be used to form
more reactive intermediates through the standard oxyhalogen
reaction R6:22
BrO3- + Br- + 2H+ f HBrO2 + HOBr
(R6)
BrO3- + CH3NH(NH2)CdS + H+ f
HBrO2 + CH3NH(NH)C-S-OH (R3)
The rate of reaction with reaction R6 as the rate-determining

13526 J. Phys. Chem., Vol. 100, No. 32, 1996
Jonnalagadda et al.
Figure 8. Linear plot of reciprocal induction time vs [BrO3-]0 (λ )
390 nm). [MTU]0 ) 0.002 M. [H+]0 ) 0.030 M. The positive intercept
on the [BrO3-]0 axis is expected, as R has to be greater than 4/3 for
any bromine to be formed.
Figure 10. (a, top) Effect of acid on the pseudo-first-order rate constant
for the formation of bromine after the induction period: Linear plot of
k′ vs [H+]0. [BrO3-]0 ) 0.10 M; [MTU]0 ) 0.001 M. (b, bottom) Effect
of [BrO3-]0 on the pseudo-first-order rate constant, k′. [MTU]0 ) 0.002
M; [H+]0 ) 0.030 M.
Figure 9. “Peacock-tail type” traces derived from variation of [MTU]0
in large excess of [BrO3]0 (R > 50). The induction period as well as
the pseudo-first-order rate constant, stays the same. The rate of
formation of Br2, however, varies. [BrO3-]0 ) 0.15 M; [H+]0 ) 0.030
M; [MTU]0 ) (a) 0.0005 M; (b) 0.001 M; (c) ) 0.0015 M; (d) 0.002
M; (e) 0.0025 M; (f) 0.003 M.
TABLE 1: Effect of Variation of [MTU]0 on the
Pseudo-First-Order Rate Constant, k′, for the Formation of
Br2 at the End of the Induction Period. In This Series
[BrO3-]0, [H+]0 . [MTU]0
[MTU]0/M
[H+]0/M
[BrO3-]0/M
kapp /s-1
step becomes
0.0005
0.001
0.0015
0.002
0.0025
0.003
0.3
0.3
0.3
0.3
0.3
0.3
0.15
0.15
0.15
0.15
0.15
0.15
0.199
0.204
0.203
0.205
0.211
0.210
-
rate ) k[BrO3 ][H+]2[Br-]
(2)
After reaction R6, the most important reaction is R5, which is
the rapid oxidation of MTU by HOBr.
Reaction R5 is the most important oxidation route of MTU
and the other intermediate sulfur species.
HOBr will react further as in reaction R5 or any R5-type
reactions (HOBr oxidation of the MTU and other intermediates)
to produce Br- and perpetuate the reaction. If reactions R5
and R7 are rapid enough, it is easy to see how eqs 1 and 2 will
become equivalent.
Further Oxidations. A mechanism that invokes bromous
acid and hypobromous acid as the reactive species requires the
oxidation processes to proceed through a series of 2-electron
The Br- formed in reaction R5 can be used in reaction R6 to
produce more HOBr and HBrO2, but in the presence of HBrO2
(from reaction R6), all the Br- will be converted to HOBr
through the rapid equilibrium, R7:23
HBrO2 + Br- + H+ h 2HOBr
(R7)

Oxyhalogen-Sulfur Chemistry
J. Phys. Chem., Vol. 100, No. 32, 1996 13527
Figure 12. Direct reaction of Br2 vs MTU monitored at 390 nm. The
reaction is so fast that it is nearly over by the time it is detected by the
stopped-flow spectrophotometer. [MTU]0 ) 0.002 M. [Br2]0 ) (a)
0.0053 M; (b) 0.0035 M; (c) 0.0018 M.
reactions. We expect the sulfur center, as well, to be oxidized
via two-electron transfer transitions. The first step in the
oxidation of the sulfur center is reaction R5 which forms the
sulfenyl acid, CH3NH(NH)C-S-OH. The sulfenyl acid can
further be oxidized to form the sulfinic acid:
HOBr + CH3NH(NH)C-SOH f
CH3NH(NH)C-SO2H + Br- + H+ (R8)
Further oxidation of the sulfinic acid yields the sulfonic acid:
HOBr + CH3NH(NH)CSO2H f
CH3NH(NH)C-SO3H + Br- + H+ (R9)
Subsequent oxidation of the sulfonic acid should concurrently
occur with the cleavage of the C-S bond to form a urea and
2-
SO4
.
HOBr + CH3NH(NH)CSO3H + H2O f
CH3NH(NH2)CdO + SO42- + Br- + 3H+ (R10)
Possible Sulfur-Sulfur Interactions. Reactions R3-R5
and R8-R10 represent oxybromine-sulfur compound interac-
tions. Some sulfur- sulfur interactions may be necessary in this
system to maintain stoichiometric consistency especially in
conditions of excess MTU. Since no colloidal sulfur was
observed when MTU is in excess, there must be some fast
rearrangements that leave the sulfur center at either oxidation
state -2 (as in MTU) or at oxidation state +6 (as in SO42-).
The other possible oxidation state, 0, will not be stable in this
environment as a sulfenyl acid. Sulfenyl acids can either
dimerize or autooxidize:
Figure 11. (a, top) Effect of [Br-]0 on the absorbance traces at 390
nm. Low Br- concentrations do not affect the rate of reaction nor the
induction time. [BrO3-]0 ) 0.15 M; [MTU]0 ) 0.002 M; [H+]0 ) 0.30
M. [Br-]0 ) (a) 0.000 01 M; (b) 0.0001 M; (c) 0.0005 M. (d) 0.001
M; Addition of Br- up to 1 × 10-5 M gave the same trace (a). These
low Br- concentrations did not affect the reaction kinetics. (b, middle)
Semilog plot of [Br-]0 vs the induction time. The data used are from
Figure 11a. (c, bottom) Semilog plot of [Br-]0 vs the pseudo-first-
order rate constant for the formation of bromine. Low [Br-]0 do not
affect the rate constant.
2CH3NH(NH)C-SOH h
CH3NH(NH2)CdS + CH3NH(NH)C-SO2H (R11)
The sulfinic acid, in the absence of further oxidant, can also

13528 J. Phys. Chem., Vol. 100, No. 32, 1996
Jonnalagadda et al.
TABLE 2: Mechanism of the Oxidation of 1-Methyl-2-thiourea by Bromate in Acidic Medium
reaction
no.
reaction
V (forward rate), V (reverse rate)
M1
M2
M3
M4
BrO3- + Br- + 2H+ h HBrO2+ HOBr
2.1[BrO3-][Br-][H+]2, 1.0 × 104[HBrO2][HOBr]
2.5 × 106[HBrO2][Br-][H+], 2.0 × 10-5[HOBr]2
8.9 × 108[HOBr][Br-][H+], 1.10 × 102[Br2]
5.3 × 101[BrO3-][SC(NH2)(NH)CH3][H+],
5.0 × 10-6[HBrO2][HOSC(NH)(NH)CH3]
1.0 × 104[HBrO2][SC(NH2)(NH)CH3]
1.0 × 104[HOBr][SC(NH2)(NH)CH3
HBrO2 + Br- + H+ h 2HOBr
HOBr + Br- + H+ h Br2 + H2O
BrO3- + SC(NH2)(NH)CH3 + H+ h HBrO2 + HOSC(NH)(NH)CH3
M5
M6
M7
M8
HBrO2 + SC(NH2)(NH)CH3 f HOSC(NH)(NH)CH3 + HOBr
HOBr + SC(NH2)(NH)CH3 f HOSC(NH)(NH)CH3 + Br- + H+
Br2 + SC(NH2)(NH)CH3 + H2O f HOSC(NH)(NH)CH3 + 2Br- + 2H+
Br2 + HOSC(NH)(NH)CH3 + H2O f HO2SC(NH)(NH)CH3 + 2Br- + 2H+
Br2 + HO2SC(NH)(NH)CH3 + H2O f HO3SC(NH)(NH)CH3 + 2Br- + 2H+
1.0 × 103[Br2][SC(NH2)(NH)CH3]
2.7 × 107[Br2][HOSC(NH)(NH)CH3]
M9
1.5 × 107[Br2][HO2SC(NH)(NH)CH3]
M10
M11
M12
M13
M14
M15
M16
M17
Br2 + HO3SC(NH)(NH)CH3 + 2H2O f SO42- + OC(NH2)(NH)CH3 + 2Br- + 4H+ 1.0 × 102[Br2][HO3SC(NH)(NH)CH3]
HOBr + HOSC(NH)(NH)CH3 f HO2SC(NH)(NH)CH3 + Br- + H+
HOBr + HO2SC(NH)(NH)CH3 f HO3SC(NH)(NH)CH3 + Br- + H+
4.0 × 104[HOBr][HOSC(NH)(NH)CH3]
5.0 × 104[HOBr][HO2SC(NH)(NH)CH3]
1.03 × 101[BrO3-] [HOSC(NH)(NH)CH3][H+]
9.0[BrO3-][HO2SC(NH)(NH)CH3][H+]
2.5 × 104[HBrO2][HOSC(NH)(NH)CH3]
1.5 × 104[HBrO2][HO2SC(NH)(NH)CH3]
HOBr + HO3SC(NH)(NH)CH3 + H2O f SO42- + OC(NH2)(NH)CH3 + Br- + H+ 1.0 × 104[HOBr][HO3SC(NH)(NH)CH3]
BrO3- + HOSC(NH)(NH)CH3 + H+ f HO2SC(NH)(NH)CH3 + HBrO2
BrO3- + HO2SC(NH)(NH)CH3 + H+ f HO3SC(NH)(NH)CH3 + HBrO2
HBrO2 + HOSC(NH)(NH)CH3 f HO2SC(NH)(NH)CH3 + HOBr
HBrO2 + HO2SC(NH)(NH)CH3 f HO3SC(NH)(NH)CH3 + HOBr
autooxidize:
and rate-determining group. Thus the global reaction dynamics
will be determined by reaction R16.
2CH3NH(NH)C-SO2H h CH3NH(NH)C-SOH +
Post-Induction Period. The rate of consumption of MTU,
the length of the induction period, and the formation of bromine
all appear to follow the same rate law, which has the term
k[BrO3-][H+]2. Thus these three events would appear to be
controlled by reaction R6. At the end of the induction period,
with all the reductant consumed, the BrO3--Br- reaction will
dominate. Bromine is formed via reaction R13, but since R13
is fast, it is not rate-determining, and the rate of Br2 formation
will be controlled by the rate of formation of HOBr (reaction
R6):
CH3NH(NH)C-SO3H (R12)
Further reactions analogous to R12 will finally produce MTU
2-
and SO4
.
The Induction Period. There is only one reaction that forms
bromine, and it is the reverse of the bromine hydrolysis reac-
tion:24
HOBr + Br- + H+ h Br2 + H2O
(R13)
-
d[Br2]/dt ) d[HOBr]/dt ) k[BrO3 ][H+]2[Br-] (3)
Reaction R13 is known to be a very fast reaction. While Br2 is
being formed through reaction R13, it will be concurrently
consumed via reaction R14,25
After considering the relative concentrations of the reactive
species,
Br2 + CH3NH(NH2)CdS + H2O f CH3NH(NH)C-
SOH + 2Br- + 2H+ (R14)
d[Br2]/dt ) kapp[Br-]
(4)
as well as in other reduction reactions that Br2 can undergo
with other sulfur compound intermediates. In general, these
reactions can be grouped as
where
-
k
app ) k[BrO3 ][H+]2
Br2 + S(X) + H2O f 2Br- + S(X+2) + 2H+
(R15)
With the amount and rate of formation of Br- being proportional
to [MTU]; Br- in eqs 3 and 4 can be replaced by MTU
([Br-]formed ) 1.33[MTU]0). The data in Figure 10a,b support
the form of eq 3.
The two-electron oxidation occurs by means of an oxygen
transfer onto the sulfur atom.
Our experimental data show that the group of reactions
indicated in reaction R15 is very fast and that Br2 will not coexist
with MTU nor with any of its intermediate oxidation products.25
Thus as long as there is reductant in the reaction solution, Br2
will not accumulate, and the reaction will remain in the
“induction period” step. Three major competing reactions are
of importance during the induction period. The first is the
general oxidation of the sulfur compound by BrO3- to produce
Br-:
Effect of Br-. Br- is a major reactant/intermediate/product
within the reaction network. Equation 3 explains why Br-
shortens the length of the induction period. Very low concen-
trations of Br- will have no effect on the rate of reaction (see
-
Figure 11b,c) because normal BrO3 solutions contain an
equilibrium amount of Br- ions that cannot be totally elimi-
nated.26 Generally, the Br- concentrations are on the order of
10-6-10-5 M. This explains why a threshold value is needed
before bromide can exert any catalytic effect. The catalytic
effect of Br- is through the enhancement of reaction R17 at
the start of the reaction. The Br2 formed in R17 will participate
in R15-type reactions to accelerate the rate of consumption of
MTU. Thus the induction period will be reduced (see Figure
11a).
BrO3- + [- - -S(-2)] + ... f Br- + SO42- + ... (R16)
Secondly, there is the BrO3--Br- reaction which produces
Br2:22
BrO3- + 5Br- + 6H+ f 3Br2 + 3H2O
(R17)
Computer Simulations
Thirdly is the general R15-type group of reactions which
consume Br2. Among these three processes, R16 is the slowest
A simplified mechanism that accounts for the observed
features of the reactions was used for the simulations. The

Oxyhalogen-Sulfur Chemistry
J. Phys. Chem., Vol. 100, No. 32, 1996 13529
-
simulations reported here are for the situation with excess BrO3
.
Excess BrO3- conditions give Br2 formation after an induction
period. To support the mechanism, the simulations needed to
correctly predict the induction time, the rate of formation of
Br2, and the mass balance in terms of the bromine and
oxybromine species. Such features are not observed in condi-
tions of excess MTU. Excess BrO3- conditions allowed us to
ignore the sulfur-sulfur reactions, which simplified the mech-
anism greatly. The simulation reactions are divided into three
groups: the oxybromine reactions, the initiation reactions, and
the oxybromine-sulfur compound reactions. These reaction
are shown in Table 2.
Oxybromine Reactions. Only three oxybromine reactions
were considered relevant for this reaction system: reactions R6,
R7, and R13 (these have been renumbered in Table 2 as
reactions M1, M2, and M3, respectively).
Reaction R6. This is the most important reaction in this
mechanism. This reaction has been extensively studied, and
the kinetics parameters are well-known. The values used for
the simulations were obtained from the FKN mechanism of Field
et al.27
Reaction R7. This reaction controls the bromous acid
concentrations. It is essential in generating the reactive species
HOBr. It is a very rapid reaction, which proceeds as written
with an equilibrium constant of 4 × 1013 M-1. The kinetics
parameters adopted for simulations were from the work of
Noyes.28
Reaction R13. The hydrolysis of Br2 was studied by Eigen
and Kustin using temperature-jump spectrophotometry. The
forward rate constant of 8 × 109 M-2 s-1 was derived from
their work.24
Initiation Reactions. These reactions were unnecessary if
one assumed a high enough initial Br- concentration in the
reaction mixture (about 10-4 M). The relevance of these
reactions was to produce enough Br- to permit the reaction to
start producing its own Br- and other oxyhalogen intermediates.
Reaction R3 is the initiator reaction when coupled with further
reduction of the HBrO2 (reactions R4 and R5). The three
initiator reactions were irrelevant when Br- was initially added
to the reaction mixture (see Figure 11a). The initiation reactions
are labeled as M4-M6 in Table 2. They also did not influence
the overall global reaction kinetics. The simulations were also
not very sensitive to the kinetics constants used for reaction
R3 when Br- was initially added as a reactant. The value of
kM4 used of 53.0 M-2 s-1 was chosen from the best simulations
fit.
Oxybromine-Sulfur Compound Reactions. For these sets
of reactions, two major oxidants were assumed: Br2 and HOBr.
HBrO2 was important as one of the initiation reactions (reaction
M5), and by assuming a rapid reaction R7, HBrO2 was also the
most important precursor to the formation of HOBr. Four
reductants were assumed: MTU and the sulfenyl, sulfinic, and
sulfonic acids. Thus eight reactions were needed to produce
an exhaustive list of oxybromine-sulfur compound reactions
(M6-M13). Since the basic assumption was that such reactions
were fast, the kinetics constants chosen did not affect the
simulations results as long as they were fast enough so as not
to be rate-determining.
Figure 13. (a, top) Comparison between experimental (solid line) and
simulations (dotted line) for the induction period and formation of Br2
in conditions of excess [BrO3-]0 using the mechanism shown in Table
2. There is good agreement in terms of the induction period, the rate
of formation of bromine, and the material balance. [BrO3-]0 ) 0.15
M; [MTU]0 ) 0.002 M; [H+]0 ) 0.30 M. (b, bottom) Simulation results
on other relevant species in the reaction. The simulations accurately
predict that the formation of bromine commences after the MTU has
been consumed (in these conditions this is about 2 s). Sulfate is formed
at about the same rate as bromine, and the sulfonic acid initially
accumulates before falling as sulfate is formed.
simulations (reactions M16, M17) allowed the simulations to
converge much faster in terms of stoichiometric consistency but
were not significant in terms of rate of reaction due to the rapid
reversibility of reaction R7.
The 17 rate equations in Table 2 were numerically integrated
using the semiimplicit Runge-Kutta method devised by Kaps
and Rentrop.29 The results of the simulations are shown in
Figure 13a,b. There appears to be very good agreement between
experimental data and the simulations. Figure 13b also proves
earlier observations that showed that sulfate formation occurs
after the induction period and at about the same rate as bromine
formation.
Additional Oxybromine-Sulfur Compound Reactions.
-
Though BrO3 is a very poor oxidizing agent compared with
Br2 and HOBr, two extra reactions were added: M14 and M15.
This was necessary because, despite its slow rate of oxidation,
BrO3- is in overwhelming excess over the reactive intermediates
and, through pure mass action kinetics, ends up being a
significant oxidation pathway. Use of HBrO2 oxidations in the

13530 J. Phys. Chem., Vol. 100, No. 32, 1996
Jonnalagadda et al.
Conclusion
References and Notes
(1) Part 19 in the series Nonlinear Dynamics in Chemistry Derived
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The mechanism reported here shows that the previous
mechanism suggested for the bromate oxidation of the unsub-
stituted thiourea was correct.15 More experimental data have
been presented here than in previous studies of oxidation
mechanisms of sulfur compounds. The data that show the
depletion of the substrate at λ ) 239 nm have been especially
useful in confirming the proposed mechanism. The reaction
with thiourea is faster than the reaction with MTU. The studies
with MTU provide experimental evidence that shows that
bromine formation takes place only after all the sulfur compound
is consumed. Bromide shows the same effect in both reactions,
i.e., a linear dependence at low concentrations, which becomes
nonlinear at high concentrations.15 A 14-step mechanism was
proposed for the bromate-thiourea reaction.15 The mechanism
proposed for the bromate-MTU system has 17-steps, which
include the formation and oxidation of the intermediate HO2-
SC(NH)NHCH3, a sulfinic acid. This is supported by previous
studies of the kinetics and mechanism of the oxidation of
aminoiminomethanesulfinic acid by oxyhalogens.25 Dispro-
portionation reactions of the oxidized sulfur compound inter-
mediates were included in the mechanism of the bromate-
thiourea system. These have been omitted from the bromate-
MTU mechanism, as new experimental data show that sulfur-
sulfur interactions are insignificant in conditions of excess
oxidant. Computer simulations of the mechanisms for both
systems show that once the sulfonic acid, HO3SC<, is formed,
its consumption is relatively slow. This is the stage where the
C-S bond is cleaved to give sulfate. The study reported in
this paper thus supports the general mechanism proposed for
the oxidation of the sulfur center in thiocarbonyl compounds
by oxyhalogens and shows that substitution at the nitrogen
affects only the rate of reaction.
(11) Capozzi, G.; Modena, G. in The Chemistry of the Thiol Group;
Patai, S., Ed.; Wiley: London, 1974; p 792.
(12) Jacobson, J. G.; Smith, L. H. Physiol. ReV. 1968, 48, 424.
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(17) Feil, D.; Loong, S. W. Acta Crystallogr. 1968, B24, 1334.
(18) Alamgir, M.; Epstein, I. R. Int. J. Chem. Kinet. 1985, 17, 429.
(19) Simoyi, R. H. J. Phys. Chem. 1986, 90, 2802.
(20) Simoyi, R. H.; Masere, J.; Muzimbaranda, C.; Manyonda, M.; Dube,
S. Int. J. Chem. Kinet. 1991, 23, 419.
(21) Brown, H. H. Anal. Chem. 1959, 31, 1844.
(22) Bray, W. C.; Liebhafsky, H. A. J. Am. Chem. Soc. 1935, 57, 51.
(23) Betts, R. H.; MacKenzie, A. N. Can. J. Chem. 1951, 29, 666.
(24) Eigen, M.; Kustin, K. J. Am. Chem. Soc. 1962, 84, 1355.
(25) Chinake, C. R.; Simoyi, R. H.; Jonnalagadda, S. B. J. Phys. Chem.
1994, 98, 545.
(26) Reagent grade Fisher sodium bromate contains 0.05% bromide by
mass (see Fisher catalog).
(27) Field, R. J.; Koros, E.; Noyes, R. M. J. Am. Chem. Soc. 1972, 94,
8649.
(28) Noyes, R. M. J. Phys. Chem. 1986, 90, 5407.
(29) Kaps, P.; Rentrop, P. Numer. Math. 1979, 23, 25.
Acknowledgment. We would like to thank the University
of Durban-Westville for granting leave to SBJ during the
summer of 1995. This work was supported by a grant from
the West Virginia-NSF EPSCoR program.
JP960259X