5124
J . Phys. Chem. 1987, 91, 5 124-5 128
Oxidatjon of Thiourea by Aqueous Bromine: Autocatalysis by Bromide' Reuben H. Simoyi2 and Irving R. Epstein* Department of Chemistry, Brandeis University, Waltham, Massachusetts 02254 (Received: February 3, 1987)
The reaction between thiourea and aqueous bromine was studied in the pH range 1.5-4. The reaction occurs in two stages: a very fast initial stage in which 1 mol of bromine is consumed for each mole of thiourea, followed by a slower second stage in which the rest of the bromine is consumed. The stoichiometry of the reaction at pH Z 2 is 4Br2 + CS(NH2), + 5H20 8Br- + CO(NH,), + S042-+ 10H'. At pH 5 2, the stoichiometry is 4Br2 + CS(NH2), + 6H20 8Br- + 2NH4+ + S042-+ C 0 2 + 8H'. The second stage of the reaction is autocatalytic in bromide. At 25.0 f 0.1 "C and ionic strength 0.2 M (NaCIO,), the rate expression is -'/jd[Br2]/dt = [CS(NH2)2][Br2](kl+ k2[Br-]), with k l = 27.8 f 0.5 M-l s-' and k2 = (3.17 f 0.3) X IO3 M-2 s-l. This reaction is explained by successive electrophilic attacks on the sulfur center by bromine.
-
-
Introduction Few kinetics studies have been performed on reactions involving sulfur-containing compounds. Reactions involving sulfur-containing compounds are generally known to be quite complex. This complexity is highlighted by free-radical mechanism^,^ autoxidation~,~ and the formation of sulfur-sulfur bonds leading to various polymeric sulfur species5 The kinetics and mechanisms of the oxidation of sulfides to sulfates, in particular, have been found to be most baffling.6 Recently, there has been an upsurge in interest in sulfur chemistry, especially in the kinetics and mechanisms of the oxidation of sulfur compounds. Much of that interest results from recent studies which have shown that chemical oscillations in a continuously stirred tank reactor (CSTR) can be obtained with sulfur-containing compounds.' Since relatively little is known of the kinetics and mechanisms of reactions of thio compounds, the sulfur-based systems are the least understood of all chemical oscillators. Field and Burger8 discovered the first halogen-free chemical oscillator: the reaction of sulfide, sulfite, methylene blue, and oxygen. Since then, other oscillators involving the reactions between hydrogen peroxide and ~ u l f i d ebromate ,~ and thiocyanate,I0 hydrogen peroxide and thiosulfate,l' and other variations have been developed. The wide variety of dynamical behavior, which includes chemical chaosI2 and "crazy clock^",'^ is comparable to that found in the better known Belousov-Zhabotinskii o~cillator.'~ Thiourea is one of the simplest of the thio compounds, yet its oxidations by iodate,I5 chlorite,16 and bromide1' have been found to give complex kinetic behavior. With iodate, the reaction displays oligooscillation, in which the concentration of iodide goes through several maxima in a single r e a ~ t i 0 n . lIn ~ a closed (batch) system, the reaction between chlorite and thiourea shows a long induction period followed by a rapid production of C102.'6 The induction period is explained by invoking a two-step process in which the thiourea reduces the chlorite to HOC1, followed by the reaction (1) Part 40 in the series Systematic Design of Chemical Oscillators. Part 39: Orbln, M.; Epstein, I. R. J . Am. Chem. SOC.1987, 109, 101-106. (2) Permanent address: Department of Chemistry, University of Zimbabwe, Mount Pleasant, Harare, Zimbabwe. (3) (a) Hashimoto, S.; Sunamoto, J. Bull. Chem. SOC.Jpn. 1966, 39, 1207-1211. (4) Wallace, T. J.; Pobiner, H.; Baron, F. A,; Schriesheim, A. J . Org. Chem. 1965, 30, 3147-3151. (5) Barnard, D. J . Chem. SOC.1957, 4675. (6) Wilson, I. R.; Harris, G. M. J . Am. Chem. SOC.1960, 32, 4515-4517. (7) OrbBn, M.: De Kepper, P.; Epstein, I. R. J . Phys. Chem. 1982, 86, 43 1-433. (8) Burger, M.; Field, R. J. Nature (London) 1984, 307, 720-721. (9) OrbBn, M.; Epstein, I. R. J . Am. Chem. SOC.1985,107,2302-2305. (IO) Simoyi, R. H . J . Phys. Chem. 1987, 91, 1557-1561. (11) Orbln, M. J . Am. Chem. SOC.1986, 108, 6893-6898. (12) Orbln, M.; Epstein, I. R. J . Phys. Chem. 1982, 86. 3907-3910 (13) Nagypll, I.; Epstein, I. R. J . Phys. Chem. 1986, 90,6285-6292. (14) Hudson, J. L.; Mankin, J. C. J . Chem. Phys. 1981, 74, 6171-6177. (15) Rabai, G.; Beck, M. T. J . Chem. Soc. 1985, 1669-1672 (16) Alamgir, M.; Epstein, I. R. Int. J . Chem. Kine!. 1985, 17, 429-439. (17) Simoyi, R . H . J . Phys. Chern. 1986, 90, 2802-2805.
0022-3654/87/2091-5 124$01.50/0
between HOC1 and chlorite to give C10,. The reaction between thiourea and excess bromate in acidic medium is also characterized by a long induction period which precedes a slow evolution of bromine." Here, too, the induction period results from a two-step process in which thiourea reduces bromate to bromide followed by the reaction of bromide with bromate to give bromine. Both the chlorite-thiourea and the bromate-thiourea reactions yield oscillatory behavior (potential of Pt electrode) in a CSTR. The bromate-thiourea oscillator differs greatly from most other bromate o~cillators.'~ The period of oscillation is very long, about 30-70 min depending on temperature, and the amplitude of oscillation is quite high (>400 mV). In excess thiourea, bromine is instantly decolorized to give bromide. It is thus tempting to assume that the origin of oscillations in the bromate-thiourea reaction lies in the [Br-]/[Br2] ratio. High bromide concentrations favor the production of bromine through the bromate-bromide reaction,18 while high bromine concentrations favor bromide production through the bromine-thiourea reaction." Under the conditions of oscillation, acidic mixtures of bromate and thiourea in a CSTR, we expect the following three reactions to be most significant: (a) the reaction of bromate with thiourea to give bromide; (b) the reaction of bromate with the bromide generated in (a) to yield bromine; and (c) the reaction of the bromine produced in (b) with thiourea to regenerate bromide. We report in this paper a kinetics study of the last of these processes, the aqueous bromine-thiourea reaction. This reaction plays a pivotal role in the bromate-thiourea oscillator. A full understanding of this reaction will be a major step toward a mechanistic characterization of the oscillator.
Experimental Section Materials. The following analytical grade chemicals were used without further purification: bromine (Aldrich), sodium perchlorate, sodium bromide, and perchloric acid, 70% (Fisher). Thiourea (Fisher) was recrystallized before use. The perchloric acid was standardized with standard N/ 10 sodium hydroxide. The aqueous bromine was standardized by reacting dilute aqueous bromine solutions with excess potassium iodide to release iodine which was titrated against standard sodium t h i o ~ u l f a t e . ~This ~ standardization was used to deduce an absorptivity coefficient for bromine at X = 390 nm of 171 mol-' dm3 cm-I. Methods. The reaction was followed by monitoring the absorbance at 390 nm on a Beckmann Model 25 spectrophotometer, the redox potential, bromide-specific electrode potential (Orion), and the pH. Redox potential measurements were performed by using a Radiometer K601 HglHg,SO41K2SO4reference electrode against a platinum electrode. The K601 reference electrode was used instead of the usual calomel reference to avoid introducing chloride ions into the reaction mixture. All reactions were performed at 25 & 0.1 O C and 0.2 M ionic strength (NaC104). (18) Vogel, A. I . Quantitative Inorganic Analysis, 3rd ed.; Wiley: New York, 1961; p 383. (19) Skoog, D. A,; West, D.M. Fundamentals of Analytical Chemistry; Holt, Rinehart and Winston: New York, 1963; p 486.
0 1987 American Chemical Society
The Journal of Physical Chemistry, Vol. 91, No. 19, 1987 5125
Oxidation of Thiourea by Aqueous Bromine Quantitative measurement of the pBr- using the bromide-specific electrode was not possible because of interference from the sulfide ions derived from thiourea, We were, however, able to make some qualitative measurements. The Davies equationZowas used to convert hydrogen ion activity (pH) to hydrogen ion concentration in those experiments in which the pH was measured. Solutions of Br3- were standardized by making a 0.001 M bromine solution in 1 M sodium bromide. According to the equilibrium Br2(aq)
+ Br- G Br3-,
Kes = 17 (ref 21)
nearly 95% of the bromine is converted to the tribromide ion. The absorbance of this solution at 390 nm was monitored, and by adjusting for the absorbance from the bromine, a value of 1006 mol-' dm3 cm-' was deduced as the absorptivity coefficient of Br3-. The bromide ion has no appreciable absorbance at 390 nm. The reactions with added sodium bromide were fast enough to be followed on a stopped-flow apparatus. The trace from the stopped flow was captured on a Biomation 610B transient recorder. Stoichiometric Determinations. Reaction mixtures used for stoichiometric determinations were allowed to sit for several hours to equilibrate before any quantitative analyses were performed. In one set of experiments, the initial thiourea concentration was held constant while the bromine concentration was varied. At the exact stoichiometric Br2:CS(NH2)2ratio of 4: 1, the absorbance of bromine at 390 nm vanishes. In excess thiourea, the remaining CS(NH2)2was determined by titration with Hg(N0,)2 in the presence of diphenylcarbazide indicator.22 Excess bromine concentrations were determined iod~metrically.'~ The sulfate ions produced were determined gravimetrically as BaS04. For reactions performed with no added acid, the acid released by the reaction was determined by titrating with 0.005 M N a O H with phenolphthalein as indicator and by calculating the hydrogen ion concentration from the pH measurements. N o special precautions were taken in the storage of the aqueous bromine. Because of the high volatility of the bromine, the solution was standardized spectrophotometrically at regular intervals of about 30 min during a set of experiments.
Results Stoichiometry. Our results indicate that the stoichiometry of the reaction under study is 4Br2 + CS(NH2)2+ 5 H 2 0 8Br- + CO(NH,), + + 10H' ( R l )
-
in the pH range 2-4. At lower pH the stoichiometry is 4Br2
+ CS(NH2)2+ 6 H 2 0 8Br-
-+
+ (NH4)2S04+ COZt: 8Hf
(R2)
Gravimetric analysis gave 97% of the sulfate ions expected from the above stoichiometry. The ratio of 4 mol of Br2 consumed for each mole of CS(NH2)2was consistently maintained. Gravimetric analysis of bromide as AgBr (after removal of S042-)also gave the above stoichiometry. N o bubbles of gas are evolved at high pH. Reactions performed in the presence of acid showed the evolution of a gas in the reaction mixture. This gas is produced by the hydrolysis of urea in strongly acidic and strongly basic solutions according to23 CO(NH2)2
+ H20
-
COZ
+ 2NH3
(R3)
The hydrolysis of urea has no effect on the 4:l ratio of bromine to thiourea. Solutions of bromine and urea can be kept indefinitely with no appreciable loss in bromine concentration except that (20) Davies, C. W. J . Chem. SOC.1938, 2093. (21) Latimer, W. M. The Oxidation States of the Elements and Their Potentids in Aqueous Solutions, 2nd ed.; Prentice-Hall: Inglewood Cliffs, 1964; p 60. (22) Yatsimirskii, K. B.; Ashtasheva, A. A. Zh. Anal. Khim. 1956, 1 1 ,
442-447.
(23) Kolthoff, I. M.; Sandell, E. G.Textbook of Quantiratiue Inorganic Analysis, 3rd ed.; Macmillan: New York, 1952; p 339.
Figure 1. Absorbance traces (A = 390 nm) for the bromine-thiourea reaction at different acid concentrations to show the effect of acid on the initial stage of the reaction. Very fast initial stage not shown; absorbance at time zero = 0.684. [CS(NH,),], = 0.002 M; [Br2I0= 0.004 M. I
12
I
2
0
0
5
10 TIME Iminutesl
15
z.loln
CSINHII;
20
Figure 2. Redox potential traces at different initial thiourea concentrations with [BrJ,, = 0.0036 M. Potential observed is essentially that of the Br,/Br- couple. When the bromine is consumed, the potential falls,
denoting the end of the reaction. expected from normal evaporation of bromine. Under the conditions of most of our experiments, stoichiometry R1 predominated. Reaction Dynamics. On mixing thiourea and bromine in neutral to slightly acidic pH, there is initially a very fast reaction in which 1 mol of bromine is consumed for each mole of thiourea. After this, the remaining bromine is consumed much more slowly (Figure 1). The redox potential remains nearly constant at about 1.05 V in the early stages of the reaction. In stoichiometric excess of bromine, the value of about 1.05 V does not change. In stoichiometric excess of thiourea, however, the system displays clock reaction characteristics in which the potential suddenly drops to about 0.6 V at the end of the reaction (Figure 2). The sharp drop in potential corresponds to the point of total consumption of bromine in Figure 1. In reactions with bromide ions initially added, the reaction displays oligooscillatory behavior, whereby one or two maxima in the absorbance at 390 nm are observed at the end of the first stage of the reaction (Figure 3). Thiourea Dependence. In a typical experiment, the initial thiourea concentration was varied while the initial bromine concentration was kept constant at 4 X M. The reaction was much faster in higher [CS(NH2),I0, as determined by how long it took for the bromine to be completely consumed. Consumption of bromine was determined by absorbance measurements and redox potential traces. Figure 2 shows three redox potential traces at different [CS(NH2),I0 values. One complicating factor was that different [CS(NH,),], values, at constant [Br,(aq)l0, gave different bromine concentrations at the beginning of the second stage of the reaction based on the 1:l mole ratio of bromine
5126
The Journal of Physical Chemistry, Vol. 91, No. 19, 1987
035t
Simoyi and Epstein
I
0251
\ 0
2
3 TIME (minutes)
4
5
consumed by thiourea in the fast initial stage of the reaction. Initial rate measurements were thus difficult to evaluate in conditions where [Br2(aq)lowas just slightly higher than the [CS(NH21210. Bromine Dependence. In these sets of experiments, the [CS(NH2)Jo was kept constant at 2.5 X M. Increasing the initial bromine concentration lengthened the duration of the reaction. A plot of reaction time vs. [Br2(aq)lohas a convex shape (Figure 4a), suggesting that some product of the reaction has a catalytic effect. The available bromine at the start of the second stage also varied with [Br2(aq)lo. Initial rates of bromine consumption at the start of the second phase were evaluated. The initial rate-[Br2(aq)lo plots were straight lines with an extrapolated [Br2(aq)loaxis intercept which was always equal to the initial thiourea concentration (Figure 4b). This intercept confirms that no second stage of the reaction will be observed when [Br2lq= [CS(NH,),],, because the thiourea will consume all the bromine in the fast initial step. Effect of Bromide. Bromide has a profound catalytic effect on the reaction. This is extremely important, since bromide is a product of the reaction. With bromide added initially, the reaction still has two stages, but the absorbance at the start of the slower phase of the reaction is very high because the Br,formed from the bromine and bromide ion has a higher absorptivity coefficient at 390 nm than bromine. The initial rate, measured as A(absorbance)/At, increases linearly with the initial bromide concentration (Figure 5a). A plot of the inverse of reaction time against initial bromide concentration gives a straight line. This implies that the rate of the bromide-catalyzed reaction dominates the other contributions to the rate. Figure 5b shows such plots. Effect of Acid. Acid has an inhibitory effect on the reaction, especially the first phase. With pH I 2.5, the initial phase does not proceed to completion and less than 1 mol of bromine is consumed for each mole of thiourea. Figure 1 shows the effects of acid on the absorbance traces, while the plot at low pH in Figure 5b illustrates the inhibitory effect of acid on the overall reaction time. Owing to the suppression of the initial phase by acid, bromine concentrations at the beginning of the second phase increase with the initial acid concentration. The initial rate of consumption of bromine thus increases with acid, though there is a quick saturation. Other Results. Bromide-specific electrode and pH measurements were also taken. Though quantitative determinations of bromide ion concentrations were not possible, the qualitative measurements showed that initially there is a very rapid increase in [Br-] corresponding to the initial phase of the reaction. After the first stage, there is a slower increase in bromide concentration to its final value. The pH-time traces were very similar to the bromide-specific electrode traces. Corresponding to the initial phase of the reaction, there is initially a rapid decrease in pH followed later by a gradual decrease in pH until the bromine is
5
4
3
Figure 3. Absorbance trace at the beginning of the second stage of the reaction showing the increase in absorbance caused by the accumulation of Brc from the reaction between bromine and bromide. [CS(NH,),], = 0.001 M, [Br,], = 0.003 M.
[Br,]
x
lo3/ M
I
b
/
I
Figure 4. (a) Plot of reaction time vs. initial bromine concentration. The shape of the plot suggests some form of autocatalysis. [CS(NH,)& = 0.0025 M. (b) Initial rate, A[Br2]/At, vs. initial bromine concentration. [CS(NH,),], = 0.0025 M; [Br-1, = 0.005 M. finally consumed. On applying the appropriate activity correctioq2O one notices that during the initial phase of the reaction 2 mol of hydrogen ions are formed for each mole of bromine and thiourea consumed. In order to test our mechanistic hypotheses (see below), several experiments were carried out in anhydrous conditions with acetonitrile as the solvent. A white precipitate was formed immediately on adding concentrated bromine to the thiourea solution. One mole of bromine reacted with 1 mol of thiourea to form this precipitate. The precipitate readily dissolved in water.
Mechanism and Rate Equation The bromide and bromine dependence results suggest two possible routes for the oxidation of thiourea in the second phase of the reaction. One route is catalyzed by bromide; the other is not. The fast initial stage involves the rapid electrophilic attack of bromine on the sulfur:
INT- 1
The Journal of Physical Chemistry, Vol. 91, No. 19, 1987 5127
Oxidation of Thiourea by Aqueous Bromine
reaction is (R4)+ (R5),in which 1 mol of bromine reacts with 1 mol of thiourea to give 2 mol each of acid and bromide ions. This concurs with our data. Further electrophilic attack by bromine should occur on the sulfur atom in INT-2. The rate-limiting step is the oxidation of sulfur from the zero oxidation state (sulfenic acid) to the +2 oxidation state (sulfinic acid). HO-S-C=NH + Brp + H 2 0
-
I
NHZ
0
II
HO-S-CZNH
+
I
2Br-
+
2H+ ( R 6 )
NH2 INT-3
Sulfur in the +2 oxidation state is quite labile and so we expect the oxidation of INT-3 to sulfate to be rapid: 0
II
HO-S-C=NH
25
I
-
I
I
0
n
1
I
I
1
I
+ 2Brp + 3 H 2 O
I
-
NH2
I
4Br-
No acid initially added 0 1 5 M acid
+
CO(NHp)2
+
6H+
+
SO,'-
(R7)
Bromide catalyzes the reaction by stabilizing the intermediate formed after electrophilic attack on INT-2. We propose an intermediate of the form INT-4:
-
20-
HO-S-C=NH
c
.-c
+
I
Brp
Br-
-
NHZ
E
I
N
z
,Br. HO-S'-'C=NH
X
I-
I Br
z
INT-4
I
W
+
15-
e / U 4
I NH,
+
Br-
H20
INT-3
The overall reaction is
:1 1
(HN)(H,N)CSOH
10
+ Br2 + Br- + H 2 0
+ 3Br- +
-
(HN)(H2N)CS02H+ 3Br-
2H+ ( R 8 )
+ 2Hf
(R9)
(-)/,)d[Br,l/dt = [CS(NH2)21 [Brzl(k1 + k,[Br-l)
(1)
which is autocatalytic in bromide. The rate law thus becomes
b
-
5 I
1
I
I
I
I
I
I
,
2
3 4 5 6 1 8 [Br-] x 10Z(MI Figure 5. (a) Rate of change of absorbance, A(absorbance)/At,against [Br-1, for the catalyzed reaction. [CS(NH,),], = 0.0025 M; [Brzlo = 0.004 M. (b) Plot of (reaction time)-' against initial bromide concentration. The inverse of the reaction time is proportional to the rate of reaction. [CS(NH,),], = 0.002 M; [Br,], = 0.0038 M.
INT-1 has other resonance structures in which the positive charge is either on the carbon or on the nitrogens. In the absence of water or any other nucleophile INT-1 is stable, with the positive charge stabilizing against further electrophilic attack. Hydrolysis of the sulfenyl halide, INT-1, yields the sulfenyl acid, INT-2:
HO-S-C=NH
I
+
2H+
+ Br-
(R5)
Initial rate measurements with [Br,], > [CS(NH,),], (low contribution from bromide) gave a value of k, = 27.8 f 0.5 M-] s-l, while in high bromide concentrations, [Br-1, >> [Br,],, [CS(NH2)2]0,the value of k2 was deduced as k2 = (3.17 f 0.3) X lo3 M-2 s-'.
Discussion Sulfur is more likely to be protonated than nitrogen when thiourea is placed in an acidic medium. The structure of thiourea nitrate crystals, for example, shows that the proton is more likely to be on the sulfur.24 Thiourea is also used as the starting material in the preparation of longer chain thiols.2s In this synthesis, thiourea is reacted with alkyl halides and the positive alkyl group attaches to the sulfur in almost 100% yield. Thus, electrophilic attack in thiourea should occur on the sulfur. From thermodynamic considerations, we expect the first phase of the reaction to be essentially irreversible toward the formation of INT-1. This suggests that acid should have no effect on the initial phase of the reaction. We found, however, that acid has a strong inhibitory effect on the initial phase (Figure la). There
NHZ INT-2
The second stage of the reaction now becomes the reaction between bromine and INT-2. The stoichiometry of 'the first stage of the
(24) Feil, D.; Loong, W. S. Acta Crystallogr.,Sect. B 1968, 824, 1334-1339. ( 2 5 ) Cossar, B. C.; Fournier, J. 0.;Fields, D. L.; Reynolds, D. D. J . Org. Chem. 1962,27, 93-95.
5128
The Journal of Physical Chemistry, Vol. 91, No. 19, 1987
Simoyi and Epstein
are two ways in which acid could have such an effect. Protonation of either or both of the nitrogen atoms would lower the potential for going from thiourea to INT- 1. The second and more likely way is that protonation of the sulfur deactivates thiourea for electrophilic attack. The protonated thiourea molecule is so stable, as evidenced by thiourea nitrate, that it can coexist with INT-I. Its attack by bromine will be much slower than attack on the unprotonated species. A similar study in which iodine oxidizes thiourea shows that iodide inhibits the reaction.ls This is explained by suggesting that I,- is less reactive than iodine. We expect the Br,- to be a much poorer electrophile than bromine as well. The equilibrium for the formation of Br,- from Br, and Br-, however, does not lie overwhelmingly to the right as in the case of The first stage of the reaction produces bromide ions. These bromide ions react with the bromine left after the first stage to produce Br3- which accumulates enough to show an increase in absorbance at the beginning of the second phase. Bromide-catalyzed brominations are common in additions of bromine across carbon-carbon double bonds.*' This type of catalysis might occur in the addition of bromine across the carbonsulfur double bond during the first stage of the reaction. We have found the first stage to be fast, and thus this catalysis cannot be observed in the Br2-CS(NH2), reaction. Bromide catalysis is observed in the second stage of the reaction during the oxidation of INT-2. Electrophilic addition is greatly enhanced if the intermediate immediately following electrophilic attack can be stabilized. Bromide, being a good nucleophile, can affect this stability by neutralizing the positive charge on the sulfur after electrophilic attack by bromine. Sulfenyl halides (e.g., INT-1) are unstable with respect to nucleophilic attack,,* and in water they hydrolyze to form sulfenyl acids:
chiometry of the initial stage have shown that with bromine the equilibrium in (R12) lies to the left and that the sulfur-sulfur bond is not favored over the sulfur-bromine bond. This also suggests that the dimerizations and disproportionations of sulfenyl halides and sulfenyl acids are not favored in the presence of bromine as the oxidizing agent. Early work on the oxidation of thiols by halogens reported sulfonic acids as the final product with a 1:3 stoichiometric ratio of thiol to halogen.33 The oxidation of thiourea by iodine15 and bromine yields a final oxidation product of sulfate with a 1:4 stoichiometric ratio. This apparent discrepancy can be explained by examining the special reactivity associated with thiourea. Urea, the oxygen analogue of thiourea, is quite unreactive. Although sulfur is not as electronegative as oxygen, the carbonsulfur double bond is much more polar than the carbon-oxygen double bond.34 The larger size of the sulfur atom militates against proper overlap of the p-orbitals that form the a-bond. Thus a large partial negative charge resides on the sulfur atom, activating it for electrophilic attack and the adjacent carbon for nucleophilic attack. This accounts for the very fast initial stage of the reaction. The two nitrogen atoms adjacent to the carbon, with their ability to stabilize a positive charge and to donate electrons, facilitate the cleavage of the carbonsulfur bond in going from the sulfonic acid to the sulfate. Although free-radical mechanisms are possible in the oxidation of thio compounds, we have proposed a mechanism here that agrees with our data without invoking radical species. We have suggested four consecutive electrophilic attacks in the transformation from thiourea to sulfate, each of these steps involving a two-electron transfer. Thus sulfur goes from an oxidation state of -2 to +6 in even steps, skipping the intermediate free-radical states. The lack of evidence for any polymeric sulfur species, even the disulfide, during the oxidation of thiourea by bromine further supports our exclusion of free-radical mechanisms. It seems likely, RSX H,O T;? RSOH H X X = halogen however, that free-radical mechanisms may be of importance in the oxidation of thiourea by bromate, since the latter species can Sulfenyl acids are very labile, and they have never been isolated.29 generate Br02'. They quickly react with sulfenyl halides or with themselves Our investigation of the bromine-thiourea reaction constitutes (disproportionation) to give thiosulfinate esters,30 which in turn easily disproportionate to give disulfides and thio~ulfinates:~~ a first step toward elucidation of the mechanism of the bromatethiourea oscillator. The bromine-thiourea system possesses (R11) 2RS-S-R RSSR + R S S 0 2 R ossufficient nonlinearity, notably a u t o ~ a t a l y s i sto , ~ generate ~ II cillatory behavior in a continuously stirred tank reactor. 8
+
+
-
-
Sulfenyl halides can also undergo dimerization to form disulfides:32 2RSX
RSSR 1- X,
(R12)
Equilibrium in (R12) lies overwhelmingly to the right with iodine,I5 and to the left with chlorine.32 Our studies based on the stoi(26) Reference 21, p 64. (27) De la Mare, P. B. D. Elecfrophilic HaloEenation; Cambridge University Press: Cambridge, 1976; p 132. (28) Kharasch, N . Organic Sulfur Compounds; Pergamon: New York, 1961: Vol. 1. D 376. (29) Capohi, G.; Modena, G. In The Chemistry of the Thiol Group, Part Two; Patai, S., Ed.; Wiley: Bristol, U.K., 1974; p 792. (30) Kice, J. L.; Cleveland, J. P. J . Am. Chem. SOC.1973, 95, 104-1 12. (31) Kice, J. L.; Venier, C. G.; Large, G. B.; Heady, L. J. Am. Chem. SOC. 1969, 91, 2028-2035.
Acknowledgment. We acknowledge helpful discussions with James Hendrickson, Kenneth Kustin, Robert Olsen, Miklds Orban, and Philip Huskey, and we thank the University of Zimbabwe for granting contact leave to R.H.S. This work was supported by Research Grants C H E 8419949 from the National Science Foundation and 2.901.1 :2789 from the University of Zimbabwe Research Board. Registry No. Bromate, 15541-45-4; thiourea, 62-56-6; bromine, 7726-95-6; bromide, 24959-67-9. (32) Reference 29, p 794. (33) Young, H. A. J . Am. Cheni. SOC.1937, 59, 811-813. (34) Janssen, M. J. In Su[fur in Organic and Inorganic Chemistry; Senning, A., Ed.; Marcel Dekker: New York Vol. 3; p 373. (35) De Kepper, P.; Epstein, I. R.; Kustin, K. J . Am. Chem. SOC.1981, 103, 6121-6127.
