Envlron. Sci. Technol. 1988, 20, 1017-1022

Kinetics of Oxidation of Dimethyl Sulfide by Hydrogen Peroxide in Acidic and Alkaline Medium

Yusuf Gbadebo Adewuyl and Gregory R. Carmlchael*

Chemical and Materials Engineering Department, University of Iowa, Iowa City, Iowa 52242

The kinetics of oxidation of dimethyl sulfide (Me$) by hydrogen peroxide has been investigated spectrophoto- metrically in acidic and neutral medium. The rate law and other data indicate the reaction is first order with respect to both H202 and MezS and is subject to catalysis by strong acids. The catalytic effect was found to be pronounced at a pH value of 1 or less. Differences in activation energies for acidic solutions (pH <7) and neutral solution (pH 7) indicate two different reaction mechanisms were operative. The rate law in strongly acidic medium was found to be -d[Me2S]/dt = (kk + kHKH[H+])[H2O2][Me2S], where k j , is the observed rate constant in uncatalyzed neutral solu- tion, the product kHKH is the catalytic constant for the acid-catalyzed reaction, KH is the equilibrium constant for the reaction of H+ with HzOz, and kH is the rate constant for the reaction of HOOH2+ with Me2S. For solutions of the same normality, the catalytic effect of HC1 was found to be double that of H2S04 at 20 "C.

Introduction Considerable quantities of gaseous discharges and

wastewater effluents containing dimethyl sulfide (Me2S) are produced by industries engaged in the manufacture of wood pulp, cellulose (by the sulfate method), and pe- troleum (1). In addition, industrial production of MezS, ai a starting product for the synthesis of dimethyl sulfoxide (Me2SO), which is a promising organic solvent, further aggravates the danger of the penetration of Me2S into open bodies of water (2). Also, the occurrence of Me2S0 in rain suggests that it is involved in the atmospheric cycle of sulfur (3). Anthropogenic sources of Me2S in natural water include activities of marine algae and bacteria (3-5).

Dimethyl sulfide has a pungent obnoxious odor and is fairly stable in water. In addition, MezS has a very low odor threshold concentration. Concentrations of a few tenths of a milligram per liter in drinking water cause noticeably disagreeable orders and taste. However, the oxidation products-dimethyl sulfoxide (Me2SO) and di- methyl sulfone (MepS02)-are relatively odorless, nonvo- latile, and water soluble. Me2S is capable of biochemical oxidation and thus increases the BOD (biochemical oxygen demand) of water. Koptyaev (6) obtained 0.3 mg/L as the threshold concentration with respect to the effect of the substance on BOD.

Oxidation of Me2S by H202 in neutral or acidic solution to Me2S0 may provide a convenient and economical me- thod for the control of malodorous conditions associated with sulfide wastes from kraft pulping processes (7, 8). With this in mind, the kinetics and mechanisms of oxi- dation of Me$ by Hz02 in aqueous solutions, and the effect of acidity have been investigated and are reported here. Hopefully, the results of this study will help in determining the optimal conditions for the control of odor and in designing of pollution control equipments for ox- idation-absorption processes.

Background In aqueous solution at ordinary temperatures, Hz02

converts MezS quantitatively to Me2S0. The stoichiom- etry is normally written as

CH3SCH3 + HzO2 - (CHJ2SO + H2O (1) At higher temperatures and excess H202, the sulfone is formed (9):

The oxidation of the sulfide to sulfoxide is more rapid than the subsequent oxidation of the sulfoxide to sulfone (10, 11). In general, the oxidation of sulfides by hydroperoxides and hydrogen peroxide are susceptible to acid catalysis (12).

Hardly any investigations of the kinetics of Me2S by Hz02 in aqueous solutions have been published in the open literature. This may be due to the limited solubility in water and the fact that the reaction is slow in the absence of a catalyst at ordinary temperatures (13). Curci et al. (14) estimated the reaction rate coefficient in water solu- tion to be 2.6 X L mol-l s-l and suggested trace metal catalysis of the reaction by sea-salt metals, but the rates are unknown.

Experimental Section Reagents and Buffers. Stock solutions of dimethyl

sulfide (Me$) and dimethyl sulfoxide (Me2SO) were prepared from ACS reagent grade obtained from Fisher and MCB Scientific Co., respectively. The dissolution of MezS was accomplished by vigorous stirring for several hours in a flask by using a magnetic stirrer before trans- ference into the thermostated reactor. The experimental reactor system has been described in detail elsewhere (15).

The experiments were conducted at pH values 7,6,4.63, 4, and 2 by using potassium phosphate monobasic-NaOH, potassium phosphate monobasic-disodium phosphate, so- dium acetate-acetic acid, potassium biphthalate, and po- tassium chloride-hydrochloric acid, respectively, a t pH values 1 and less by using appropriate amount of HCl (6 N) and HzS04 (6 M) solutions, and with unbuffered dis- tilled water. The buffers for the pH values 7,4, and 6 were concentrates, and the appropriate amounts were added to the reactor (with the equilibrated solution of Me2S) at the start of the experiments. The buffer for pH values 2 and 4.63 were already in the required solution form, and the Me2S was dissolved directly in it.

A volume of H202 (50% solution, Fisher) calculated to give [H2O2], was added to the reactor at the initiation of the reaction.

Analytical Procedures. Spectrophotometric mea- surements were made with the Lambda 3 UV-visible double-beam scanning spectrophotometer (C618-0900, Perkin-Elmer) and Coleman cells with an optical path of 1 cm. Fisher Accumet selective ion analyzer equipped with Ag/AgCl pH (Corning 476022) electrodes was used to check the pH of the reaction solutions.

Both Me2S and MezSO have ultraviolet absorption of significant intensity in the 200-220-nm region. Figure 1 illustrates the molar absorbance coefficients of Me2S, Me2S0, and H202 in distilled water in the wavelength range of interest. The curves were obtained from prepared standard solutions of these compounds. It is apparent that Me2S0 has absorption maximum at 205 nm, and the molar absorptivities of H202 are relatively low in this region.

CH3SCH3 + 2H202 - (CH3)2S02 + H2O (2)

0013-936X/86/0920-1017$U1.50/0 0 1986 American Chemical Society Environ. Sci. Technol., Vol. 20, No. 10, 1986 1017

1600 1 l v

190 200 210 220 230 WAVELENGTH ( n m )

Figure 1. Molar absorptivities of Me& Me,SO, and H,O,.

2.5 1 - 205 nm 4 215 nm -+ 225 nm

I 0 40 80 1 eo 160 200

TIME ( m i n )

Figure 2. Absorbance as a function of time for Me,S oxidation at three wavelengths and pH 6, T = 25 O C , [Me,S], = 1.44 X M, and [HPOP]o = 1.955 X M.

Dialkyl sulfones [e.g., (CH3),SO2] are generally transparent throughout the ultraviolet region except in strongly alka- line solutions when absorption due to anion formation is observed (13).

For a mixture of independently absorbing substances, the total absorbance at a given wavelength is

(3)

where C1, C,, ... CN are the concentrations of the substances having absorption coefficients El , E2, ... EN and 1 is the path length. The components of the mixture can be an- alyzed if absorbances are measured at wavelengths where the absorption coefficients of the species in question are sufficiently different and the corresponding simultaneous equations solved. In this work, the concentrations of Me2S and Me2S0 at a given sampling time were obtained by measuring absorbances at 205 and 215 nm where absorp- tion coefficients of the species are significantly different and relatively large compared with that of H202 (see Figure 1). This allowed the absorbance of H20z to be ignored and the number of simultaneous equations to be solved reduced to two. Figure 2 illustrates typical absorbance vs. time behavior of a Me2S + H202 system at 25 O C and pH 6 and with [Me.$], = 1.44 X M. By use of E205 = 1610 and E215 = 1200 for Me2S and E205 = 960 and E,,, = 530 for Me2S0, the concentrations of these compounds in solution calculated at various sam-

1018 Environ. Sci. Technol., Vol. 20, No. IO, 1986

A, = (E1C1 + E2C2 + ... + E N C N ) l

M and [H2O2lO = 1.95 X

z P

e 0.5 - !- z W 0 z 0 0

G

m

X

-0- DMS - DMSO - DMS +DMSO

I I I I I I

0 4 0 00 120 160 200 2 4 0 TIME (min)

Figure 3. Concentrations of Me$ and Me2S0 as a function of time during oxidation for pH 6, T = 25 OC, [Me,S], = 1.44 X M, and [H20,], = 1.955 X M.

Table I. Summary of Kinetic Data for Me2S Oxidation

[MezSI, WZOZI, PH T, k X lo2, X103 M X103 M (or acidity) "C M-'s-l

1.444 1.955 6 25 4.9 1.289 1.955 unbuff (H20) 25 2.6 1.450 1.955 7 25 1.6 1.302 1.955 1 (0.1 N HC1) 20 4.3 1.450 1.955 1 (0.1 N HCl) 25 6.1 1.440 1.955 2 20 3.5 1.451 1.955 6 20 3.4

1.407 1.955 7 20 1.4 1.408 1.955 2 30 6.3 1.334 1.955 1 (0.1 N HC1) 30 8.1 1.447 1.955 unbuff (HzO) 30 3.8 1.450 1.955 6 30 6.5 1.355 1.955 7 30 1.7

1.423 1.955 unbuff (H,O) 20 2.2

1.386 1.955 0.2 M H2S04 20 5.7 1.306 1.955 0.1 M H2S04 20 3.5 1.285 1.955 0.15 M HzSO4 20 4.5 1.293 1.955 0.05 M HzS04 20 2.4 1.305 1.955 0.2 N HCI 20 7.4

pling times are plotted in Figure 3. As can be seen from the smooth curve obtained, the sum of [Me2S] and [Me2SO] at a given time is approximately equal to [Me2SI0. Thus, ignoring the absorbance due, H202 has no significant effect on the results and is therefore justified.

Spectrophotometric analyses were impossible at pH values 4 and 4.63 because of interferences from compo- nents of buffers at wavelengths of interest.

Results and Discussion Determination of k. In order to determine the rate

constants for the oxidation of MepS by HzOz in acidic and neutral medium, a series of experiments at different [Me2SIo, pH, and temperatures were conducted (experi- mental conditions are summarized in Table I). In all runs, the concentrations of H202 used was kept close to the amount required stoichiometrically. Typical absorbance vs. time data is shown in Figure 2. The kinetic rate con- stants can be determined by the following analysis.

0.20

0.15

," 0.10

s -

W

0.05

0.0

- pH 1 (0.1 N HCI), [DMS], = 1.356 x 10-3M

-c- p H 6, [ DMS], = 1.445 X 10-3M

-0- unbuffered, [DMS],. 1 .289x10-3M - p H 7, [DMS], = 1.450 x 10m3M

J

0 40 SO 120 160 200 240 280 TIME ( m i n )

Figure 4. Secondorder lots of runs at various pHs, T = 25 O C , and

Consider eq 1 and let [MezS] = A, [H202] = B, and

A + B - X (4)

A = [A], - X ( 5 )

B = [B], - X (6)

dx/dt = k([A]o - x)([B]o - X ) (7)

Integrating eq 7 by partial fractions, substituting for x from 5 and 6, and rearranging give

[H,O,], = 1.955 X 10- f M.

[MezSO] = X , then k

where

A plot of In ([B]/ [A]) vs. time of experimental data should give a straight line with slope of ([B], - [A],)k. However, in this work, data for [B] is not available and so an al- ternative solution is desirable. Introducing the dimen- sionless quantities

(9)

(10)

X z = - (fraction Me2S reacted)

[AI,

(ratio of [H20,], to [Me,S],) [BI, [AI,

Q = -

eq 7 becomes

(11) dz dt - = [A]ok(l - z)(Q - Z )

and integration of eq 11 gives

8 - 2 Q(1 - 2)

In - = [A],(Q - 1)kt

0.20

0.1 5

0 0.10 E!

W s

0.05

0

-c pH 1 (0.1 N HC!), [DMS], = 1.302 X

-E- p H 6, [DMS],= 1.451 x

-c- unbuffered, [DMS], = 1.423X - pH 7, [DMS], = 1.407 X

40 80 120 160 200 240 280 TIME (min)

Figure 5. Second-order plots of runs at various pHs, and T = 20 OC, and [H,O,],, = 1.955 X M.

The same result is obtained by substituting eq 9 and 10 directly into eq 8. A plot of log [(Q - z) /Q(l- z)] VS. time should give a straight line with slope [AlO(Q - l)k/2.303, from which k is obtained.

The results for various experimental conditions are plotted in Figures 4-7 and the values of the rate constants obtained are reported in Table 1. As can be observed, the rates are dependent on the acidity of the solution and the temperature. The rate constants are lowest for the neutral solutions (pH = 7), irrespective of temperature and tend to increase with decrease in pH. For example, at 20 "C, k is 1.4 X 2.2 X and 3.4 X L/(mol.s) for solution of pH 7, unbuffered system, and solution of pH 6, respectively. The pH of the unbuffered system was lower than 7. This is expected since H202 acts as a weak acid in the presence of water. The reaction H202 + H20 F! H30+ + 02H- occurs to a small extent, and the lowering of pH is due to the hydronium ion (16, 17).

The increase in oxidation rates with decrease in pH is even more vivid in the presence of HC1 and H2S04 (i.e., with pH 1 and below). For instance, in the presence of 0.2 N HC1, the rate at 20 OC is over 5 times the corresponding value in neutral solution (pH 7), and with 0.2 M HzSO4, it is about 4 times even though the later is more acidic. The lower rates observed in H2S04 solutions (compared to those in HC1 solutions) may be due to the formation of peroxomonosulfuric acid (18), which uses up part of the available H20z according to the reaction

K=l H202 + HzS04 HOOSOZOH + HZO (13)

H2SO6 is an inorganic peroxide formed from H202 by the replacement of one hydrogen atom by the HOSO2- group. In strong peroxide solutions (go%), replacement of two hydrogen atoms is possible, leading to the formation of peroxodisulfuric acid (H2S208). H2S05 is moderately soluble in water and fairly stable in solution (13).

Environ. Sci. Technol., Vol. 20, No. 10, 1986 1019

0.20

0.15

0.10 0

P c3

41

0.5

0.0

- pH 1 (0.1 N HCI ), [DMS], = 1.334 x

-0- pH 6, [ DMS], = 1.450 x M -a- unbuffered,[DMS], = 1.447 x 10-3M - pH 7, [DMS], =, 1.355 X 10-3M

M

, I I I I I I

0 40 SO 120 160 200 240 280 TIME (rnin)

M. Flgure 6. Second-order plots of runs at various pHs, and T = 30 OC, and [H202]0 = 1.955 X

0.2

0.15

0.1 0 0 -1

0.05

0

--c 0.2 M H2S04, [DMS],= 1.386 x 10-3M

+ 0.1 M HzS04, [DMS] ,= 1.306 x IOm3M

0.05 M H2S04 [DMS],: 1.293 x 10-3M

TIME (rnin)

Figure 7. Secondorder plots of runs at various H,SO, concentrations, T = 20 OC, and [H,0210 = 1.955 X M.

Figure 8 illustrates the Arrhenius plots for k at various pHs. From the curves, activation energies (EA) of 10.94, 10.76, 11.73, and 3.41 kcal/mol are obtained for reaction solutions of pH 1, unbuffered, pH 6, and pH 7, respec- tively. The values for EA are almost identical for all the reaction systems except for the neutral solution. This

1020 Environ. Sci. Technol., Vol. 20, No. 10, 1986

8.0 7.0

6.0

5.0

4.0

3.0

c .- E ' 2.0 0 - E \

2 + .- - . 1.0 2

(B) neutral medium (uncatalyzed) 4.0

3.0

C .- E l u) 0)

0

\

- E 2.0

f 5 .- - r

1 .o

0 I I I I

0.05 0.1 1.5 2.0

[H,S04] mole/liter

Figure 9. Effect of sulfuric acid on the rate of oxidation of Me,S by HzOz at 20 OC.

tained for HC1 catalysis using the data a t 0.1 (Le., k = 2.6 M-l min-l) and at 0.2 N (i.e., k = 4.4 M-l min-l) as

kHCl = 0.8 + 18.1[HC1] (15) Similar relationships have been observed by other in-

vestigators for other reactions (21-24). Overberger and Cummins (22), in the oxidation of p,p’-dichlorobenzyl sulfide by hydrogen peroxide in the presence of H2S04, obtained 18 L mol-’ min-l as the catalytic constant for the acid-catalyzed reaction. Such catalysis is known to be of the specific acid type since general acids (such as acetic acid) often do not show catalysis with hydrogen peroxide oxidations (19). As Bateman and Hargrave (25)) observed in the oxidation of cyclohexyl methyl sulfide by hydro- peroxides in alcohols, the catalytic power of the acid is determined by its strength relative to the solvent. Ross (21), in the oxidation of thiodiglycol with Hz02, found the rate constant a t pH 1 (0.1 M HC1) to be twice the value at pH 4.7 (buffer solution 0.5 M in both sodium acetate and acetic acid) even though the latter has higher ionic strength.

Mechanisms. Mechanisms that are consistent with observed kinetic results and account for the catalytic be- havior of H+ ion are as follows: (A) acidic medium (catalyzed)

K H H+ + H202 HOOH2+ (16)

HOOH2+ + (CH3)zS - I (17) kH

where I is the intermediate: CH3 0 H

:s: - 0’ H I

CH3 H

I 5 H30’ + (CH&SO (18)

where I’ is the intermediate HCjH

n b o ~ t

CHS 6 CHS

I’ .-% HOH + HOH + (CH3),S0 (20)

Equation 16 involves the addition of a proton to HzOz to form H302+ in a rapid equilibrium. The equilibrium constant for this reaction has been estimated to be of the order of M (26). The effect of acid on the oxidation rate could be explained by supposing that H302+ (the conjugate acid of H202) can donate OH+ ion to a nucleo- phile faster and easier than the neutral HOOH in the rate-determining step (eq 17). Only in strongly acid SO- lution (pH G l ) is there enough of this species present to make a difference in rate. It is faster because the other product of its oxidation is not the OH- ion, involving a separation of charge in the critical step of the reaction, but a neutral water molecule. In the oxidation of thiocyanate ion by hydrogen peroxide, Wilson and Harris (27) observed the order of reactivity of the electrophilic substitution of HzOz to be H3OZ+ > Hz02 > HO,. The production of HOC (H20z + OH- s HO, + HzO) is pronounced only in strong alkaline medium, pH >11 (28).

Equation 19, the rate-controlling step for the uncata- lyzed reaction, represents a termolecular displacement with water, the electrophilic agent, and sulfide, the nucleophilic agent. The hydrogen peroxide could be presumed to be largely solvated prior to oxidation. Overberger and Cum- mins (22) observed that the uncatalyzed oxidation of sulfides by H20z was promoted in solvents capable of hydrogen bonding with H202, and they also found alcohols and water to be equally effective in solvating hydrogen peroxide. It is generally true that reactions producing products more polar than the reactants are favored in polar solvents (29). However, solvation of sulfoxide by solvent results in a negative charge displacement toward the solvation site with reduction in the nucleophilicity of the sulfur atom and decreased oxidation of sulfoxide to sulfone.

The overall rate expressions can be obtained by applying steady-state approximations to the intermediate I and I‘ (i.e., dI/dt = dI’/dt = 0):

-d [ Me,S] = kH[HOOH2+] [MezS] dt

and

for the cases of acid-catalyzed and uncatalyzed reactions, respectively. The overall rate law over a range of pH values can thus be described by the relation

where k = kfH + kHK~[H+] and hi, is the observed rate constant in uncatalyzed neutral solution and the product kHKH is the catalytic constant for the acid-catalyzed re- action. The existence of two independent rate laws is a characteristic of H202 reactions with electron pair donors including the thiosulfate, chloride, and hypochloride and

1021 Environ. Sci. Technol., Vol. 20, No. 10, 1986

1 I I I I I I I

PH 0 1 2 3 4 5 6 7

Figure 10. Plot of rate constant vs. pH at 20 O C .

the bromide and hypobromide (20).

Summary and Conclusions Oxidation of dimethyl sulfide by hydrogen peroxide is

bimolecular (unit order with respect to H20z and Me$) and is subject to acid catalysis. The observed rates are found to be pH dependent. As shown in Figure 10, the rate coefficient is fairly constant between pH 2 and pH 6, increases drastically at pH 1 and below, and decreases substantially at pH 7. The presence of 0.1 N HC1 solution increased the observed rate constant by a factor of 3-5 times the value in neutral solution depending on the tem- perature. The observed catalysis is not simply due to an increase in ionic strength (21). For solutions of the same normality, the catalytic power of HCl was found to be double that of H2S04 at 20 "C. The low catalytic ability observed with H2S04 solutions compared to HCl was partly attributed to the depletion of H202 in solution due to a secondary reaction with H2S04 to form peroxornonosulfuric acid (H,SO,) which is fairly stable in solution.

Two different reaction mechanisms were found to be operative as indicated by the differences in activation energies. For acidic solutions (pH <7), the average acti- vation energy was found to be about 11.14 kcal/mol whereas EA for reaction in neutral solution (pH 7 ) was only 3.41 kcal/mol. Observed rate constants in acidic solutions were found to be double with 10 "C increase in tempera- ture. Reactions in neutral medium were far less sensitive to temperature changes.

On the time scale of cloud processes, the oxidation of MezS by Hz02 may be considered slow. For instance, at pH 6 and temperature 25 OC, the conversion of Me2S to Me2S0 after 150 min is only 50% (see Figure 3). The small reaction rate suggests that it is unlikely that the aqueous oxidation of MezS by H20z can account for the presence of MezSO in rain and snow observed by Andreae (3, 5) . However, the oxidation of MezS by H20z in collected acidic rainwater samples may be appreciable on time scales of hours to a few days. Furthermore, the rate coefficients (especially in acidic medium) are large enough to provide efficient oxidation of Me2S to MezSO in other aqueous systems (e.g., lakes, reservoirs, sewage). Typical wastes- treams from kraft mills (8) contain sulfur dioxide (SO,), hydrogen sulfide (H,S), mercaptans (RSH), dimethyl

sulfides (RSR), and dimethyl disulfides (RSSR). Hydrogen peroxide will oxidize SO2 to sulfate (30), HzS to sulfate and sulfur (311, RSH and RSSR to sulfonic acid (RS03H) and sulfate (32,331, and RSR to sulfoxides and sulfones. The products of oxidation are all odorless, and hence, HzOz may provide economic means for odor and wastewater quality control.

Registry No. Me#, 75-18-3; HzO,, 7722-84-1; HC1,7647-01-0; HZSOI, 7664-93-9.

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Received for review October 21, 1985. Accepted April 21, 1986.

1022 Environ. Sci. Technol., Vol. 20, No. 10, 1986