Kinetics of Hydrolysis and Oxidation of Carbon Disulfide by Hydrogen Peroxide in Alkaline Medium and Application to Carbonyl Sulfide

Environ. Sci. Technol. 1987, 21, 170-177

Kinetics of Hydrolysis and Oxidation of Carbon Disulfide by Hydrogen Peroxide in Alkaline Medium and Application to Carbonyl Sulfide

Yusuf Gbadebo Adewuyl and Gregory R. Carmichael*

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

Kinetic studies of the oxidation of carbon disulfide by hydrogen peroxide in alkaline medium were made spec- trophotometrically. The reaction of CS2 with OH- ion was found to be rate controlling and proceeded by the formation of a dithiocarbonate complex. The major reaction product was sulfate with sulfur occurring as colloidal suspensions only at pH values less than 8. The formation of sulfate increased exponentially with time and was also found to be dependent on the rate of hydrolysis of CS2. In addition, the production of sulfate showed large induction periods, suggesting either a complex mechanism or formation by secondary reactions. The results obtained for carbon disulfide were extended to carbonyl sulfide (OCS) oxidation in alkaline solutions. The removal of OCS (acid gas) from mixtures of gases by alkaline liquid ab- sorbents (e.g., NaOH) and oxidation of subsequent solutions to sulfate is an important industrial practice.

Introduction Anthropogenic sources of CS2 in the aqueous environ-

ments include wastewater effluents from factories engaged in the manufacture of artificial leather, viscose, rayon, and other synthetic fibers. In addition, CS2 has wide industrial applications as an excellent solvent for waxes, resins, rubber, sulfur, and other substances. CS2 also naturally occurs in coal tar and crude petroleum. Recent experiments indicate that the oxidative decomposition of pyrite (FeS,) by oxygen in aqueous environments may be a source of CS2 and OCS (I).

Carbon disulfide is a poisonous and volatile liquid with a pungent smell. Apart from its toxicity to animals and aquatic organisms (2), CS2 can hydrolyze to H2S and other reduced sulfur species in aqueous environment. The hydrolysis products are malodorous and corrosive and may also lead to deposition of sulfur and disruption of biological waste treatment processes. In view of these facts, the kinetics, mechanisms, and stoichiometry of the oxidation of CS2 by H202 in alkaline medium have been investigated and are reported here. It is hoped that the results of this study will enhance our understanding of the chemical removal of CS2 and other reduced sulfur species from natural water and wastewater.

Background CS2 is quantitatively oxidized to sulfuric acid by H202

in alkaline solutions (3). The stoichiometry may be assumed to be CS2 + 8H202 + OH- +

HC03- + 2HS04- + 2H+ + 6H20 (1)

While the kinetics of oxidation of H2S by H202 in acidic solutions have been reported in the open literature ( 4 , 5 ) , such studies involving CS2 have been limited only to its hydrolysis in strong alkaline solutions (6-9).

Generally, the stoichiometry for the hydrolysis of CS2 in strong alkaline solutions can be written as

(2) cs2 + 60H- - 2s'- + C03'- + 3H20

with a secondary reaction (S2- + CS2 - CSS2-) occurring

only at pH >13 (10, 11). However, the hydrolysis has been found to involve the formation of an unstable intermediate dithiocarbonate (DTC), which can exist as CS20H- or CS202- (12). The intermediate can disproportionate in one or more rapid steps into carbonate and sulfide. The sequence of reactions depends on the pH of the solution. With pH >11, Philipp et al. (II) and Hovenkamp (7) obtained

CS, = CSzOH- C_+_ C S z 0 2 - 5 on- on- n

(very rapid)

With pH <11, the sequence given by Vermaas (13) is the most likely alternative:

HCO3- + HS- ( 4 )

Philipp (6) obtained the activation energy (EA) between temperatures of 20 and 30 "C for the decomposition of CS2 to be 10.4 kcal/mol a t pH 9 and 20.7 kcal/mol at pH between 11.3 and 12.6. Cherkasskaya et al. (9) also obtained EA in the same temperature range to be 20.3 for CS2 in 1 N NaOH solution. The difference in EA may be at- tributed to the existence of two possible mechanisms in the pH range of interest.

Hovenkamp ( 7 ) studied the hydrolysis in strong alkaline solutions (NaOH > 0.1 M) and found the rate-determining step to be first order with respect to both CS2 and OH- ion concentratiorls. He also concluded from his results that dithiocarbonate existed as CS202- in the [OH-] range studied. The DTC was found to absorb significantly from 190 to 310 nm with a maximum at 272 nm. With a 0.2 N NaOH solution, a molar absorptivity (E) of 10 500 L mol-l cm-l was obtained at 272 nm and 25 OC. The value of E

was also found to be dependent on OH- ion concentrations. At 250 nm, the molar absorptivity was found to be only 3250 L mol-' cm-'.

Experimental Section Reactor System. A schematic diagram of the experi-

mental setup is shown in Figure 1. The reaction vessel was fabricated from Plexiglas tubing. An air-tight seal was maintained around the piston, electrodes, thermometer, and the cylindrical glass frit by the insertion of O-rings. The displacement piston allowed the addition of buffer solutions and the oxidants without introduction of air. The glass frit allowed purging of the system with nitrogen to get rid of dissolved oxygen. Constant temperature in the reaction vessel was maintained externally by a Lauda thermostat equipped with a circulator system and a temperature controller. Samples to be analyzed were drawn from the reaction vessel through the sampling port with a 12-mm hypodermic syringe. Mixing was initiated by a magnetic stirrer. The reaction vessel had a total internal volume of 3.8 L.

M) were prepared by first adding the appropriate amount Reagents. Stock CS2 solutions (4.5 X lo4 to 1.5 X

170 Environ. Sci. Technol., Vol. 21, No. 2, 1987 0013-936X/87/0921-0170$01.50/0 0 1987 American Chemical Society

Figure 1. Experimental reactor system: (1) thermostated glass reaction vessel; (2) magnetic stirrer; (3) Teflon-covered magnet; (4) Plexiglass lid; (5) thermometer; (6) sampling port; (7) coarse glass frit for gas in; (8) air vent for gas out: (9) glass and reference electrodes; (10) temperature-controlled circulator system (water in and out); (1 1) Accumet selective analyzer; (1 2) Accumet electrode switch.

of Fisher reagent-grade CS2 (0.1-0.5 mL) to 1 L of distilled water in flat-bottomed flask accompanied by vigorous stirring with a magnetic stirrer. The solution was then transferred to the reaction vessel and diluted further with distilled water to the required concentration. Insulators were wrapped around the vessel, and mixing was further initiated. The solution was left overnight to attain dy- namic equilibrium. An appropriate volume of H202 (Fisher certified, stabilized, 50% solution) calculated to give [H202], was added directly to the reactor a t the start of the experiment.

Stock sulfide solutions were prepared from ACS reagent-grade Na2S*9H20. The oxidized surface coating on the crystals was first removed by scratching with a stirring rod. Other reagents included barium chloride crystals (20-30 mesh) and anhydrous Na2S04 used for sulfate calibration curves.

Buffer Solutions. Experiments presented here were conducted at pH values of 11, 10,9, 8, 7.41, and 7.0 and were buffered with NaOH-glycol-NaC1, boric acid-KC1- NaOH, borax (Na2B40,.10H20), sodium phosphate-potassium phosphate, potassium phosphate monobasic-sodium phosphate dibasic, and potassium phosphate monobasic-Na0H buffer systems, respectively. The buffers for the pH values of 11, 10, and 7 were concentrates, and the appropriate amounts were added to the reactor at the start of the experiments. The buffers for the pH values of 9,8, and 7.41 consisted of dry salta and hydrion capsules. The required amounta were dissolved in 500 mL of distilled water before introduction into the reactor. All buffers were ACS reagent grade obtained from Fisher Scientific Co.

Analytical Procedures. Two major methods were explored in following the course of oxidation of CSp by H202-standard spectrophotometric techniques (14) and potentiometric methods involving the use of selective ion

0.0 210 220 230 240 250 260 270 200 290 300 310 320

WAVELENGTH lnmm)

Figure 2. Molar absorptivities of HS-, H20,, and CS,.

eleectrodes. The spectrophotometric measurements were done with the Lambda 3b UV-visible (190-900-nm) dou- ble-beam scanning spectrophotometer (C618-0900, Per- kin-Elmer) equipped with Model RlOO Microprocessor strip recorders and Coleman cells with an optical path of 1 cm and provided with a lid. Fisher Accumet selective ion analyzer (Model 750) equipped with electrode switch (Model 753) and Ag/AgCl glass junction reference (Corn- ing 476067) and Ag/Ag2S (Corning 476129) and Ag/AgCl internal pH (Corning 476022) electrodes were used for the potentiometric measurements.

Earlier efforts to use a sulfide selective electrode (Ag+/S2-) to follow the course of reaction were unsuccessful because of uncertainties due to drifting problems. Moreover, the electrode responds only to S2- ion activity and not to the bulk of sulfide in solution. However, in the pH region of interest (7-ll), the HS- will appear to be the predominant species in aqueous solutions (15,16). Thus, the use of electrodes was limited mainly to checking of pH values.

The ion HS- exhibits strong absorption in the UV region from 190 to 260 nm with a maximum molar absorptivity of 7.8 X lo3 1 mol-l cm-l at 230 nm (17,18). As mentioned earlier, the intermediate DTC absorbs in the region 190-310 nm with a maximum at 272 nm (7). The possible oxidation products, (S032-, HS03-, SO:-, HSOc) do not absorb appreciably in the UV region 250-280 nm (19,20). Figure 2 shows the molar absorptivities of CS2, H202, and HS- in the wavelength range of interest. The curve for HS- was obtained from standard solutions of Na2S.9H20. Treiber et al. (21) obtained maximum absorption at 315 nm for CS2 in water. Hence, to determine the values of DTC and HS-, spectrophotometric measurements were made at 272 and 250 nm at each sampling time. The absorbance at 272 nm was due mainly to the dithiocarbonate (DTC) intermediate.

The turbidimetric method was used to measure sulfate produced from oxidation of CS2. This method is based upon the fact that barium sulfate tends to precipitate in a colloidal form and that this tendency is enhanced in the presence of NaC1-HC1 solution containing glycerol and other organic compounds (22-24). This method was cho- sen because (1) it is fast and simple, (2) temperature variations of 10-15 deg do not cause appreciable error, and (3) high concentration samples, which will otherwise give absorbance readings exceeding instrument limits, can be diluted and the resulting concentrations multiplied by the appropriate dilution factor. The precipitate formed absorbs strongly from 380 to 430 nm. Measurementa were made at 420 nm. Sulfate concentrations were obtained from a standard calibration curve (Figure 3). The curve

171 Environ. Sci. Technol., Vol. 21, No. 2, 1987

0.6 c

4 0.3

0.6 0.7 k

\ 0 1 2 3

CONCENTRATION (MI x io4 Figure 3. Standard caiibration curve for turbidimetrlc sulfate determination obtained at 420 nm and T = 25 f 1 OC.

E 0.3 N

Q 0.2

T = 20' C

0.1 I I I

T TIME ( m i n ) 5 10 15 20

0.6 0.7 k

0.15 1 I 1 I I I I

10 20 30 40 50 60 70 TIME ( m i n )

Figure 4. Formation and decay of DTC monltored at 272 nm for [CS2l0 = 1.372 X lo3 M, [H,O,], = 7.41 X io-', pH 9, and T = 20 OC.

is linear, following the Beer-Lambert law between 1.2 X

Experimental Procedures. The reaction vessel was washed with detergents and rinsed with distilled water. The vessel was then cleaned with a 50% solution of " 0 3 (to ensure removal of any traces of sulfur oxidation products and other impurities that might have been present on the surfaces of the vessel) and rinsed again with distilled water. The stock solution of CS2 was then in- troduced into the reactor and equilibrated as described earlier. Distilled water used in making solutions was either boiled or purged with N2 prior to use.

Prior to the start of each experiment, initial concentration of CS2, pH, and temperature of the reaction solution were checked. The reaction was initiated by adding the appropriate amount of buffer followed immediately by the addition of HzOz. The timer was simultaneously started. Time zero was taken as the time just before H2Oz was added. Aliquots for analysis of DTC, HS-, and sulfate


and 5.0 X M.

T.25" C

0.1 I I I I

0.8 T TIME (min) 0.7

0.6

5 10 15 20

0.5 (u 0.4 IC cu ;3 0.3

n 8

0' 0.2

0.1 I I I I 1 1

10 20 30 40 50 60 TIME (rnin)

Figure 5. Formation and decay of DTC monitored at 272 nm for [CSi], = 1.372 X M, [H2O2I0 = 7.41 X lo-* M, pH 9, and T = 25 C.

were withdrawn at periodic intervals through the sampling port. Temperature, pH, and S2- readings were also taken at these intervals. The reactions were continuously monitored for 8 h and the final readings taken after 24 h.

Results and Discussion Kinetic Studies. In order to determine the rate

equations for the oxidation of CS2 with H20z in alkaline medium, a series of experiments at different [CS2l0, pH, and temperatures were performed. In all runs, the concentrations of Hz02 used were far in excess of the amount required stoichiometrically. Thus, D,, which is the asymptotic absorbance at infinite time, was taken to be due to excess H202 and was considered the base-line absorption.

Typical experimental results are shown in Figures 6, 7, 10, 11, and 12. The absorbance data show a rapid increase followed by a slow decay. This behavior is indicative of a series reaction of the type A - B - C (25-27). In addition, the absorbances at 250 nm (curve B, Figure 7) are greater than those at 272 nm. Since the molar absorptivities of DTC are greater than those of the HS- ion (7) and since DTC absorbs more at 272 nm than 250 nm, these absorbance observations indicate that HS- is formed at a faster rate than DTC. The sulfate concentration data show a characteristic growth curve with the asymptotic value of twice the initial CS2 concentration. Moreover, the sulfate concentration data show a marked induction period, indicating that sulfate is formed by secondary reactions.

An overall reaction sequence consistent with the experimental data is

1.0 0.9 0.8 0.7 0.6

0.5

0.4

0.3

(u r; 0.2 c 0

8 D l

0' 0.1

0.09 0.08 0.07 0.06

0.05

0.04

0.7 0.6 0.5 0.4

0.3

0.2

0.1 0.09 0.08 -- 0.01

'c 0.06 E 0.05 x 0.04

0.03

0.02

.- -

go1 0. 09 0.008 0.007 0.006 0.005 0.004

-e- k,, SLOPE ~ 0 . 6 1 0 4 I

- - - - -

-

- - - - - - - - -

- - - - - - -

- p H = 9

+ pH = 10 -CI- p H = l l

1.0

r

c .- E 1

X

o.l t

0.03 0 10 20 30 40 50 60

TIME (min)

Flgure 6. Formation and decay of DTC monitored at 272 nm and different pHs for [CS2l0 = 4.573 X M, [H202]0 = 4.94 X IO-' M, and T = 20 OC.

0.01 L I I

10-5 10-4 10-3

Flgure 8. Effect of [OH-] on reaction rates for [CS,] = 4.573 X lo-' M, [H202]0 = 4.94 X lo-' M, and T = 20 OC.

[OH-] moles/!

B

- p H = 1 1

* pH=10

a

0.5

0.4 E

N F N

0 0.3

c

8 n , 0.2

TIME (rnin)

Flgure 7. Formation and decay of DTC (A) and DTC + HS- (B) monitored at 272 and 250 nm, respectively, for different pHs and [CS2l0 = 4.573 X IO-' M, [H202]0 = 4.94 X lo-' M, and T = 10 OC.

where MTC is the monothiocarbonate intermediate CS- 02H-. [MTC is also the intermediate in the reaction of OCS with OH- ion (28).] Pseudo-first-order conditions were ensured by holding [OH-] constant with buffers and using excess H20z (5-15 times the amount required stoichiometrically).

2.34 kcal

1.96 kcal

0.003 * 3.2 3.3 3.4 3.5 3.6 3.7

1 x 1 0 3 T ( O K - ' )

Figure 9, Effect of temperature on reaction rates for [CS2l0 = 1.372 X M, [H,02]o = 7.41 X M, and pH 9.

Determination of kl and k2. Consider the consecutive reactions A L% B h as in eq 5. The rate expressions for A (CS,) and B (DTC) are

[A] = (7)

Environ. Scl. Technoi., Vol. 21, No. 2, 1987 173

Integrating eq 8 with [B] = 0 at t = 0 gives

kl[A1o (e-klt - e-kzt) [B] = - kz - kl

The absorbance at 272 nm at a given time can be written

(10) By substituting eq 7 and 9 into eq 10 and simplifying, the net absorbance is given by

Dt = EA[A]~ + EB[B]~ + D,

D = ( ye -k l t + (11)

where D, - D, is denoted by D and

Taking the logarithm of both sides of eq 11 and differentiating with respect to time, we have

-(akle-klt + flk2e-k2t) S = (14)

( Y e - k ~ t + @e-&

where S denotes d(1og D)/dt. A semilogarithmic plot of Dt - D, vs. time thus consists

of two added straight line segments as shown in Figures 4 and 5. The segments can be resolved to provide values for the respective k's provided they are not too close to- gether. The kinetic situation A - B - C admits two possible considerations: the rapid formation of a weakly absorbing intermediate or the slow formation of a strongly absorbing intermediate (25, 29). As mentioned earlier, previous studies indicate the latter (i.e., k2 >> k,) is ap- plicable in these studies (7). Thus, as t - m, e-k2t will be much smaller than e-klt and may be so neglected. In eq 14, S - -kl as t -+ m, and also eq 11 reduces to

D = (Ye-klt (15) Hence, the long-time portion of Figures 4 and 5 (lower curves) will have slopes of -kl, and careful extrapolation of the lines to zero gives log,, a, With (Y known, a new difference is computed from eq 11:

= ae-klt - D (16)

The quantity AZl2 denotes the difference between the data at short times and the extrapolation of the long-time linear portion. A semilogarithmic plot of A212 vs. time affords k, (see Figures 4 and 5, upper curves). Figures 4 and 5 also illustrate the fact that for k, >> kl the concentration of DTC will pass through a maximum. The time for this maximum concentration is obtained by differentiating eq 9 and setting the result to zero:

'

Given kl from plots of eq 15, k2 can also be obtained directly from eq 17.

A summary of k, and kl determined experimentally is shown in Table I. Both kl and k2 are found to increase with increases in temperature, pH values, and initial CS2 concentrations. However, values reported by other in- vestigators (6-9,30) could not be compared directly with those obtained in this work because they used strong


- I O O C 1 0 50 100 150

TIME l m l n l

Flgure 10. Sulfate production as a function of time for different temperatures and [CS,], = 1.372 X M, [H,Oplo = 7.41 X lo-* M, and pH 9.

-0- pH 9 y -i=- pH 8

0 50 100 TIME (rn in)

Figure 11. Sulfate production as a function of time for different pHs and [CS,], = 1.372 X M, [H,O,]o = 7.41 X lo-' M, and T = 25 "C.

4x10.'

L 0 50 100 150

TIME l m i n l

Flgure 12. Sulfate production as a functlon of time for different pHs and [CS,l0 = 2.228 X M, [H,02]o = 7.41 X lo-* M, and T = 10 "C.

NaOH solutions and their medium was unbuffered. As shown in Figure 8, a log-log plot of both kl and kz

shows first-order dependency with respect to OH- ion concentration. The curve for k2 gives a slope of 0.61, and that for kl gives 0.32. Figure 9 illustrates the Arrhenius plot (31) for both kl and k2. From the kl curve, the activation energy for the formation of DTC was found to be 12.00 kcal/mol. Philipp (6), for a pH of 9, obtained a value of 10.4 kcal/mol for EA (temperature between 20 and 30 "C). Also from k,, a value for E A of 12.3 kcal/mol is obtained for the decomposition of DTC.

Determination of Sulfate Production Rate. Con- sider the reaction scheme given in eq 5

Table I. Summary of Kinetic Data for CSz Oxidation

I c s 2 lo P xi03 M

1.372 2.228 2.228 2.228 1.372 1.372 1.372 1.372 0.457 0.457 0.457 0.457 0.457 0.457

[HzOzlot X102 M

7.41 7.41 7.41 7.41 7.41 7.41 7.41 7.41 4.94 4.94 4.94 4.94 4.94 4.94

PH

9 9 9 8 8 9 9 9

11 10 9

11 10 9

T, OC

20 20 10 20 25 25 15 10 20 20 20 10 10 10

k1, x106 s-1

22.8 26.2 79.3

10.4 51.3 14.8 11.0 86.7 40.0 20.1 46.0 16.3 12.1

5.96

26.3 4.57 45.8 13.2 34.4 24.1 2.08 40.2 10.7 18.9 13.1

444.2 97.7 26.7

156.8 26.8 9.05

3 x lo-' I

- = - p H = 9

il- pH =IO -p- p H = 8

0 50 100 150 TIME (min)

Figure 13. Sulfate production as a function of tlme for different pHs and [CS,], = 2.228 X 10-~ M, [H,Op]o = 7.41 X IO-, M, and r = 20 OC.

and assume that the reaction of OH- ion with CS2 is rate-controlling, then the sulfate production rate can be written

where kls is the pseudo-first-order rate constant for sulfate production). From stoichiometry

d[CSz] d[HS04-] -2- - - dt d t

Comparison of eq 18-20 shows that k1s = 2k1 (21)

The rates of sulfate production were obtained from smooth curves of [S(VI)] vs. time data (Figures 10-13), and the semilogarithmic plots of (1/ [CSzI0)(d[S(VI)]/dt) vs. time are shown in Figure 14. For the cases presented in Figure 14, the slopes of the straight lines (i.e., kls in Table I) were found to be twice the values of their corresponding k l values as indicated in Table I and hence confirm the relatiohship in eq 21. The intercepts of these lines on the ordinate were also found to be approximately twice the values of kls. For example, for [CS2l0 = 1.37 X T = 20 OC, and pH 9 (with S(V1) production shown in Figure 111, the slope is 0.027 and the intercept is 0.051 (obtained from least-squares analysis). For the same conditions above but T = 25 OC, the slope is 0.064 and the intercept is 0.145. Similarly, for the case of [CSzl0 = 1.37 X T = 25 "C, and pH 8, the slope is 0.012 and the intercept is 0.021. Thus, the sulfate production rate increases with increase in pH and temperature just as the rate of for-

0.05

0.04

0.03

I

r

,E 0.02 E Y

0.01 0 0.009 N 0.008

0.006

0.005

0.004

-20" C, pH.9

+- 25' C, pH =9

0.003 7 0 20 40 60 80 100

TIME (rnin)

Figure 14. Normalized sulfate production rate as a function of time at different pHs and temperatures for [CS,], = 1.372 X M and [H,O,], = 7.41 X lo-' M.

Table 11. Sulfate Production after 24 h

[CS~IO, [HzOzlo, [S(VI)IT, ?& of xi03 M xi02 M PH T, o c xi03 M [SI,

1.372 7.41 9 20 2.365 86.2 1.372 7.41 9 25 2.270 82.9 1.372 7.41 8 25 2.365 86.2 2.228 7.41 11 10 4.320 96.9 2.228 7.41 10 10 4.430 99.4 2.228 7.41 11 20 4.140 92.9 2.228 7.41 9 10 4.240 95.2 2.228 7.41 8 20 4.130 92.7 0.915 4.94 11 10 1.850 101.1

mation and decomposition of DTC.

rate can be described in terms of [CS2l0 as From the results discussed above, the sulfate production

where kls = 2k1. The total sulfate produced after 24 h (when reactions

are assumed complete) for some experimental runs was measured and is given in Table 11. The final amounts of sulfate were found always to be twice the initial concentration of CSz. This confirms the stoichiometric rela-

175 Environ. Sci. Technoi., Voi. 21, No. 2, 1987

tionship given in eq 1. The percent conversion to sulfate based on the recoverable total sulfur, [SI, (which equals 2[CSzl0), is also computed and reported in Table 11. It can also be observed that pH plays an important role in the amount of conversion to sulfate. Conversions were on the average over 95% with pHs of 10 and 11 irrespective of temperature and [CSzl0. However, the percent conversion decreases with pH. This might be due to the appearance of colloidal sulfur observed at low pH values. In fact, the reactions could not be monitored at pH 7.41 and 7 because they were too slow and also the formation of whitish turbid suspensions made spectrophotometric readings erratic. The formation of sulfur from oxidation of sulfide by Hz02 in neutral and acidic solutions has been reported in the literature (4 ,5,32) . Another important observation made in the course of the reaction (and shared by these inves- tigators) was the fact that production of sulfates in basic solutions was always accompanied by a slight decrease in the pH of the resulting solutions. Hoffman (4 ) , in the oxidation of H2S with H202, also obtained a high sulfate yield in basic solutions. With [S2-l0 = 7.6 X [H2OZ], = 7.6 X T = 25 "C, and pH 8.5, SO:- was 99% of the recoverable total sulfur. He also found that a 4-fold excess of peroxide was needed for the reaction to go to comple- tion.

Mechanisms. A mechanism that describes the oxidation of CS2 by Hz02 under the conditions used in these experiments should (1) be consistent with previously determined rate equations, (2) identify the reaction products in the rate expression, (3) predict the overall stoichiometric requirements of the reaction, and (4) account for the catalytic behavior of the OH- ion. A mechanism that fulfills all of the above requirements in basic solution (as already discussed) can be written in the following manner:

k l CS2 + OH CS20H-

CS20H- + OH- 2 CS02H- + HS-

CSOZH- + OH- -% CO3H- + HS-

HS- + H202 -% HSOH + OH-

HSOH + H202 5 S(OH)2 + HzO

S(OH)2 + H202 2 S02.H20 + H2O k7

k-7 SOyH20 HSOf + H+

HS03- + H202 HS04- + H20

Repeating the steps in eq 25-29 (2 mol of HS--mol of CS2), the overall reaction predicted in eq 1 is obtained.

Equation 23 can be assumed to be rate-determining. The assumption appears valid if one compares the observed rate constants (Table I) with the measured rates for the other steps. Equations 24,25, 27, and 28 all involve unstable intermediates and are generally expected to be fast. The monothiocarbonate intermediate (CSOZH-) formed in eq 24 is by far less stable than the dithiocarbonate (6, 7). In equation 26, HS- acts as a nucleophile in an attack on H20z, which then undergoes a heterolytic breakdown with hydroxide as a leaving group. According to Hoffman (4) , k4 is 29.0 M-l min-l (at 25 "C and pH 5.05) and increases significantly with an increase in pH. For example, in the oxidation of H2S by H2O2 at 25 "C, Hoffman (4) also found that pseudo-first-order rate con-

176 Envlron. Sci. Technol., Vol. 21, No. 2, 1987

stant k , ~ (rate = k o ~ [ S z - 1 ~ , where [S2-]~ = [HzSI + [HS-I + [S2-l, kobsd = k,[Hz021[H+], and kc is the overall rate constant) to be 2.609 at pH 8.1 and only 0.021 at pH 5.05. Equation 29 is an equilibrium reaction that is nearly in- stantaneously established. At 20 "C and p = 0.1 (ionic

determined by ultrasonic absorption (33). The same in- vestigators obtained a value of 7.2 X lo2 M-' s-l for k, (at pH 6.4, 12 "C, and p = 1). The reaction (eq 30) proceeds via a nucleophilic displacement by H202 on HS03- to form peroxomonosulfurous acid (02-SOOH), which is an unstable intermediate (34, 35).

At low pHs, the formation of sulfur follows a different mechanism. The HSOH intermediate formed in eq 26, instead of reacting with Hz02 to form the sulfoxylic acid intermediate [S(OH),], reacts with HS- to form HS2- and water. The formation of higher polysulfides by the reaction of HSOH continues in subsequent steps until HSg- is formed. The final step involves the formation of the HS- and c-SS by an intermolecular displacement of HS; (4,36). The c-SS is the only thermodynamically stable form of sulfur a t STP (16,37). Teder et al. (38) also suggested a complex reaction between sulfite and sulfide leading to the formation of sulfur a t pH values under 8.

Summary and Conclusions The oxidation of CS2 by H202 to sulfate in alkaline

medium (pH 111) is controlled by the hydrolysis step. The reaction of CS2 with OH- ion is the rate-determining step and proceeds by the formation of a complex dithiocarbonate (CS20H-). The linear increase as a function of OH- ion concentration for the pseudo-first-order rate constants for both the formation and decomposition of this complex (as in Figure 8) leads to the concept of bimolecular course for both steps.

With pH 8 and above, the predominant product of oxidation is sulfate, and in this case an 8-fold excess of peroxide will be required for the reaction to go to com- pletion. The formation rate of the sulfate is linear with respect to the CSz concentration and is dependent on the rates of hydrolysis. Thus, an increase in pH results in increased sulfate production rates. The productions of sulfate also show large induction periods suggesting either a complex mechanism or formation by secondary reactions. Moreover, the fact that orders in the rate laws are less than the stoichiometric numbers suggests the formation of an intermediate in the rate-determining step.

With pH 7.41 or 7, formation of colloidal sulfur was observed. The resulting milkiness due to the suspensions of sulfur, coupled with the fact that the hydrolysis step was very slow (as observed from measured absorbance changes), made spectrophotometric analysis impossible in this pH range.

It is evident from the results summarized above that, under the proper conditions, H202 appears to be an effective reagent for the control of unpleasant conditions resulting from CS2 in aqueous systems. The advantages of using H202 over other potential oxidants (e.g., 0 2 , O?, C12) are its decomposition products (02, H20) are nontoxic (39), it is a liquid and can conveniently be applied as needed, it ensures complete oxidation to nontoxic sulfate with elimination of pungent smell, and it is economical compared with other treatment systems. Furthermore, H202 is an effective oxidant for other reduced sulfur com-

strength), k, = 3.4 X lo6 s-l and k-, = 2 X 10, M-l s-l a S

pounds (40). Amlication to OCS Oxidation. Carbonyl sulfide is

vicir&y poisonous, more so than H2S and CS2, and it is dangerous because it is odorless, tasteless, and colorless. It constitutes a major contaminant in oil and synthetic gas

industries. It is important to remove OCS from process gas streams because it is very corrosive (especially in humid conditions) and is also capable of poisoning catalyst in subsequent operations. Absorption and oxidation of the gas in alkaline medium to sulfate by H202 may provide an economical method.

The rate data for the alkaline hydrolysis of OCS are well established in the open literature. While the rate is faster than that of CS,, it is slower compared with rates in eq 26-30. Philipp et al. (28) observed that the rate-determining step in the alkaline hydrolysis of OCS was the formation of the monothiocarbonate intermediate (MTC) according to a bimolecular course:

OCS + OH- 2 CS02H-

At a pH value of 10, pseudo-first-order rate constants for this reaction were given as 0.02,0.042, and 0.086 m i d for 10, 20, 30 bC, respectively. An activation energy of 13.2 kcal/mol was obtained for the formation of MTC and 23.1 kcal/mol for the faster decomposition step. Philipp et al. (28) also observed a linear increase of k9 as a function of OH- concentration and hence a bimolecular rate coefficient of 4.8 X 1O1O ea3/T M-l s-l. Sharma (41), in experiments conducted in a stirred tank and wetted wall column, obtained a value of 14.0 M-l & for the bimolecular rate in 1 M NaOH solution (pH 14) and at 25 OC. Sharma and Danckwerts (42) also obtained a value 12 M-l for the rate constant of 25 “C.

The decomposition of MTC to sulfide and subsequent oxidation to sulfate should follow the same steps as in eq 25-30. Thus, the overall stoichiometry may be written as OCS + OH- + 411202

HCO, + HSO, + H+ + 3H20 (32) From the data given above, the overall kinetics in any buffer solution should be governed by the rate of hydrolysis of OCS. Hence, in analogy to CS2 oxidation, sulfate production rate may be assumed to be given by

(33)

where kg’ is the pseudo-first-order rate constant; kg’ = k,[ OH-].

Registry No. CS2, 75-15-0; HzOz, 7722-84-1.

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Received for review October 25, 1985. Revised manuscript received July 15, 1986. Accepted September 25, 1986.

Envlron. Sci. Technol., Vol. 21, No. 2, 1987 177

Documents

Kinetics of Hydrolysis and Oxidation of Carbon Disulfide by Hydrogen Peroxide in Alkaline Medium and Application to Carbonyl Sulfide