Sonochemical Oxidation of Carbon Disulfide in Aqueous Solutions

Greensboro, North Carolina 27411. The kinetics of sonochemical oxidation of CS2 and the effects of process parameters (e.g., concentration, pH, temper...
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Ind. Eng. Chem. Res. 2002, 41, 4957-4964

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Sonochemical Oxidation of Carbon Disulfide in Aqueous Solutions: Reaction Kinetics and Pathways Yusuf G. Adewuyi* and Collins Appaw Department of Chemical Engineering, North Carolina A and T State University, Greensboro, North Carolina 27411

The kinetics of sonochemical oxidation of CS2 and the effects of process parameters (e.g., concentration, pH, temperature, ultrasonic intensity, irradiation medium, dissolved gas and time, etc.) on the degradation rates and product distributions were studied in a batch reactor at 20 kHz. Sonochemical oxidation was found to decrease with an increase in the solution temperature, and the reaction rate order was dependent on the temperature (T) range studied: zero-order at T g 10 °C [with rate constants of (0.66-3.68) × 10-5 M min-1] and first-order at T e 5 °C (with rate constants of 0.037-0.266 min-1). From Arrhenius law, k ) A exp(EA/RT), the activation energy, EA, for the zero-order degradation of CS2 in the presence of air as the irradiating gas was found to be 7.2 kJ/mol at the higher temperatures compared to 28.7 kJ/mol for the firstorder degradation at the lower temperatures. Sonochemical oxidation pathways leading to sulfate formation are discussed. The results of this study suggest that the ultrasonic degradation of CS2 might provide an environmentally conscious method for the control of this hazardous pollutant in industrial wastewater. Introduction Biogenic sulfur (or reduced sulfur) compounds are the main causes of odor in natural waters and process wastewaters, and control of their emissions is a key environmental concern in the purification of natural gas, viscose rayon manufacture, tanneries, and the kraft pulp and petroleum refining industries.1 These compounds include carbon disulfide (CS2) and its hydrolysis products in an aquatic environment, hydrogen sulfide (H2S), and carbonyl sulfide (OCS).2,3 Carbon disulfide (CS2) is a poisonous, volatile, and pungent-smelling liquid with wide industrial applications as an excellent solvent. It is also toxic to animals and aquatic organisms, and its aqueous hydrolysis products, H2S and OCS, are also malodorous and corrosive. When released into the atmosphere, CS2 is mostly converted to SO2 and OCS in the lower atmosphere, with the latter being one of the greenhouse gases with a lifetime of more than 1 year in the atmosphere.4 It is also classified as a hazardous air pollutant under Title III of the 1990 Clean Air Act Amendment of the United States.5 The development of cost-effective and environmentally conscious technologies for its control is therefore desirable and of an increased interest.4-6 The kinetics and mechanisms of oxidation of carbon disulfide in aqueous solutions by hydrogen peroxide and other chemical oxidants (e.g., O3, Cl2) at different pHs and temperatures were studied in detail by previous investigators.1-3 However, today’s treatment processes must meet more stringent requirements of being environmentally responsible, and innovative technologies such as sonochemical oxidation, requiring little to no chemical additions for effective remediation, are more desirable.7,8 Sonochemical techniques utilize ultrasound to produce an oxidative environment via acoustic cavitation due to the formation and subsequent collapse of * Corresponding author. Phone: (336) 334-7564. Fax: (336) 334-7904. E-mail: [email protected].

microbubbles from acoustical wave-induced compression/rarefaction. The collapse of these bubbles leads to localized transient high temperatures (g5000 K) and pressures (g1000 atm), resulting in the generation of highly reactive species including hydroxyl (•OH), hydrogen (H•), and hydroperoxyl (HO2•) radicals and hydrogen peroxide. The reactive species are capable of initiating or promoting many reduction-oxidation reactions. Adewuyi8 provides a comprehensive review of the fundamentals of environmental sonochemistry discussing applications to contaminant remediation and challenges for scale-up and commercialization. Entezari et al.9 studied the sonochemical degradation of pure liquid carbon disulfide and the effects of frequency, temperature, intensity, and gases on the rate of its dissociation. They found that ultrasonic irradiation of the CS2 liquid at 20 kHz resulted in the formation of a heterogeneous mixture of black particles (amorphous carbon) in a yellow solution (monoclinic sulfur). We investigated the kinetics of the sonochemical oxidation of unbuffered aqueous carbon disulfide earlier at 20 kHz and 20 °C in a batch reactor.10 With an initial CS2 concentration of (13.2-13.6) × 10-4 M, we found the reaction rate to be zero-order, and the rate constant for the degradation at 20 °C and 14 W (11.04 W/m2) in air was 21.1 µM/min compared with 46.7 µM/min at 50 W (39.47 W/m2). The formation of sulfate as the main reaction product was enhanced in the presence of hydrogen peroxide but inhibited in the presence of 1-butanol. To further understand the kinetics of the sonochemical oxidation of CS2 and the reaction pathways, we have extended this study to include reactions at different temperatures (1-50 °C), unbuffered solutions, and solutions with pHs of 8-11. The results of these studies are reported here. Experimental Procedures The experimental setup for this study consisted of a 20 kHz sonifer capable of a maximum power output of

10.1021/ie020069a CCC: $22.00 © 2002 American Chemical Society Published on Web 08/31/2002

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400 W (Branson model 450 sonifer), jacketed glass reactor, sound abatement enclosure box, and circulatory water bath described elsewhere in detail.10 Kinetic runs were carried out by ultrasonically irradiating the resulting solutions for a desired length of time, using intensities of 14-50 W, and at temperatures of 1-50 °C. At the start of each experiment, a prepared stock solution (50 mL) of CS2 with known concentration (e.g., 0.02, 0.03, or 0.04 M) was poured into the reactor. The reactor was immediately sealed to prevent any contaminants from volatizing from the stock solution. A gas dispersion tube was inserted in the reactor below the surface of the solution. An irradiating gas, air, helium, argon, or nitrous oxide was bubbled into the sample for a period of 30 min to saturate the solution. The bubbling process resulted in the reduction of the initial concentrations of the stock CS2 solutions of 0.02, 0.03, and 0.04 M to typically (6.4-7.0) × 10-4, 10.5 × 10-4, and (13.2-13.9) × 10-4 M, respectively, due to volatilization. These resulting concentrations were recorded as the initial CS2 concentrations at the start of ultrasonic irradiations. Degradation of the gas-saturated CS2 solutions was monitored at 314 nm at different time intervals using a Beckman DU-7000 spectrophotometer. The products of the reaction were analyzed using a DIONEX ion chromatograph (with an IonPac AS11 2 mm analytical column and an EG 40 eluent generator using 1 M potassium hydroxide as the eluent) at the end of each irradiation period. The experimental procedure was repeated for different sets of process parameters (e.g., pH, temperature, solute concentration, power or intensity, dissolved gases). Experiments presented here were conducted with unbuffered solutions, and at pH values of 11, 10, 9, and 8 buffered with NaOH-glycol-NaCl, potassium carbonate-potassium borate-KOH, boric acid-KCl-NaOH, and potassium phosphate monobasic-NaOH buffer systems, respectively. These buffers were concentrates, and appropriate amounts were added to Milli-Q water and CS2 dissolved in the solutions at the start of the experiments. The solution pH was measured with a Fisher Scientific pH/ion conductivity meter (model 50). All of the buffer solutions were ACS reagent grade obtained from Fisher Scientific Co. Ultrapure-grade air, nitrous oxide (N2O), helium, or argon used as the nucleating or saturating gas was obtained from Air Products Co., Ltd. Results and Discussion In our earlier study, which was carried out in a batch reactor at a frequency of 20 kHz and 20 °C, the sonochemical oxidation of CS2 was shown to follow zeroorder with the production of sulfate as the product.10 To further understand the kinetics of the sonochemical oxidation of CS2 and the reaction pathways, experiments were conducted at different process parameters (e.g., intensities, initial aqueous CS2 concentrations, irradiation time, and medium), temperatures ranging from 1 to 50 °C, unbuffered solutions, and solutions with pHs of 8-11, and the reaction products were analyzed. The experimental conditions and results are summarized in Table 1 and Figures 1-11. Table 1 also includes two data points at 20 °C (P ) 14 W) with air as the irradiating gas: [CS2]0 ) 6.90 × 10-4 and 13.36 × 10-4 from our earlier study10 duplicated here for the benefit of comparison. As shown in Figure 1, CS2 is effectively degraded by ultrasonic irradiation, with the rate of degradation

Figure 1. Concentration of a CS2 solution as a function of time during ultrasonic irradiation in the presence of air at different temperatures {[CS2]0 ) (6.84-7.00) × 10-4 M, P ) 14 W}. Table 1. Summary of the Results for the Degradation of CS2 [CS2]0 (×104 M)

temp (°C)

gas present

power (W)

rate constant (k × 105 M/min)

R2

7.00 6.95 6.90 6.94 6.90 13.36 13.33 13.34 13.33 13.34 13.34 13.79 13.45

A. Zero-Order Degradation of CS2 10 air 14 2.43 15 air 14 2.31 20 air 14 2.27 30 air 14 2.01 50 air 14 1.67 20 air 14 2.11 20 helium 14 3.68 50 air 30 1.81 50 air 50 3.46 50 helium 14 1.89 50 air 14 1.17 50 N2O 14 0.86 50 argon 14 0.66

0.9993 0.9953 0.9981 0.9938 0.9929 0.9966 0.9883 0.9914 0.9991 0.9896 0.9892 0.9797 0.9813

6.91 6.84 6.93 10.52 13.38 13.40 13.37 13.18 13.33 13.43 13.90

B. First-Order Degradation of CS2 1 air 14 2.08 3 air 14 1.91 5 air 14 1.73 5 air 14 1.07 1 air 14 0.70 1 air 30 1.24 1 air 50 2.66 5 air 14 0.59 5 air 30 0.80 5 air 50 1.30 5 N2O 14 0.37

0.9905 0.9795 0.9840 0.9853 0.9856 0.9925 0.9789 0.9712 0.9910 0.9644 0.9808

increasing with a decrease in the temperature. The dependency of rate constants on the temperature is shown on Arrhenius plots (Figures 9 and 10) and will be discussed later. In general, a lowering of the ambient temperature increases the rate of sonochemical reactions unlike most other chemical reacting systems. However, this will depend on whether the cavitation occurring is transient (vaporous with little or no dissolved gas present) or more resonant (stable or gaseous with mainly dissolved gas and some vapor). The maximum temperatures and pressures attained during the collapse of transient cavitation bubbles are predicted by Noltingk and Neppiras from approximate solutions of the now so-called Rayleigh-Plesset-Noltingk-Neppi-

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Figure 4. Zero- and first-order kinetic degradations of CS2 at 5 and 20 °C with air as the irradiating gas {[CS2]0 ) (13.2-13.4) × 10-4 M, P ) 14 W}.

Figure 2. Concentration of a CS2 solution as a function of time during ultrasonic irradiation in the presence of different gases {[CS2]0 ) (13.2-13.9) × 10-4 M, T ) 5 °C, P ) 14 W}.

Figure 5. Zero- and first-order kinetic degradations of CS2 at 5 and 50 °C with N2O as the irradiating gas {[CS2]0 ) (13.8-13.9) × 10-4 M, P ) 14 W}.

Figure 3. Concentration of a CS2 solution as a function of time during ultrasonic irradiation in the presence of different gases {[CS2]0 ) (13.4-13.8) × 10-4 M, T ) 50 °C, P ) 14 W}.

ras-Poritsky (RPNNP) bubble dynamic equation assuming adiabatic collapse and are given in eqs 1 and 2,11-14 where T0 ) temperature of the bulk solution, Pv

Tmax ) T0

[

Pmax ) Pv

[

]

Pa(γ - 1) Pv

]

Pa(γ - 1) Pv

(1)

γ/(γ-1)

(2)

) pressure in the bubble at its maximum size, Pa ) pressure in the bubble at the moment of transient collapse (i.e., acoustic pressure), and γ ) polytropic

Figure 6. Zero-order kinetic degradations of CS2 at the higher temperatures (T ) 10, 15, 20, 30, and 50 °C) with air as the irradiating gas {[CS2]0 ) (6.9-7.0) × 10-4 M, P ) 14 W}.

index or heat capacity ratio of the cavity medium.8,11 For vaporous cavitation, increasing the ambient temperature, T0, will raise the equilibrium vapor pressure of the medium and so lower both Tmax and Pmax. As the temperature of the liquid, T0, is increased, its vapor

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Figure 7. First-order kinetic degradations of CS2 at different initial concentrations, [CS2]0 (T ) 5 °C, P ) 14 W).

Figure 9. Arrhenius plot of CS2 degradation rate constants at the higher temperatures (T ) 10, 15, 20, 30, and 50 °C) with air as the irradiating gas {[CS2]0 ) (6.9-7.0) × 10-4 M, P ) 14 W}.

Figure 10. Arrhenius plot of CS2 degradation rate constants at the lower temperature (T ) 1, 3, and 5 °C) with air as the irradiating gas {[CS2]0 ) (6.84- 6.93) × 10-4 M, P ) 14 W}. Figure 8. First-order kinetic degradation of CS2 at the lower temperatures (T ) 1, 3, and 5 °C) with air as the irradiating gas {[CS2]0 ) (6.90-7.00) × 10-4 M, P ) 14 W}.

pressure, Pv, is also increased, but much more dramatically than the temperature. The vapor, which enters the bubble during its formation, cushions the collapse of the bubble and dampens the dissipation of ultrasonic energy. This predicts that, as the bulk temperature increases, the temperature of the “hot spot” formed by the collapsing cavity decreases, resulting in a decrease in the reaction rate as indicated by the results of this study. It should be noted that, in certain reaction systems, more favorable results are attained at an optimum reaction temperature. In such systems, an increase in the ambient temperature is found to increase the kinetic reaction to a point before the cushioning effect of the vapor in the bubble begins to dominate, resulting in a decrease in the reaction rate upon further temperature increase.11 In the investigation of the degradation of thymine, Sehgal and Wang15 also found that the rate

may even reach a plateau with an increase in the temperature before decreasing upon a further increase in temperature. Studies also indicate that the nature of cavitation prevailing is determined to a large extent by the ultrasonic frequency of the system. Entezari and Kruss16,17 studied the sonochemical reaction rate of iodide oxidation at different temperatures (0-50 °C) and at 20 and 900 kHz. At 20 kHz the reaction rate decreased with an increase in the temperature at all power levels, but at 900 kHz the rate showed a maximum at a temperature intermediate between 5 and 50 °C, dependent on the power level. They observed that the noise caused by the collapse of the cavitation bubbles was very noticeable at 20 kHz compared with that at 900 kHz and concluded that transient cavitation was more likely occurring at the 20 kHz frequency and resonant cavitation at the higher frequency of 900 kHz.17 The effect of temperature on degradation rates observed at 20 kHz in this study is consistent with the results of our studies. However, using a model which linked bubble dynamics with the production of free

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Figure 11. Sulfate formation as a function of time for solutions of different initial temperatures, T, in the presence of air as the irradiating gas {[CS2]0 ) (13.60-13.8) × 10-4 M, pH ) 11, P ) 50 W}.

radicals, Sochard et al.18,19 predicted a maximum for both •OH and H• radicals with an increase in the liquid bulk temperature. They attributed these results to two competitive but opposite effects: (1) an increasing liquid temperature leading to an increased amount of water vapor in the bubble, which can promote formation of free radicals from dissociation of water molecules; (2) increasing liquid temperature leading to less violent collapse, which can result in lower internal temperatures at the end of the collapse phase. As shown in Table 1, the sonochemical reaction rates are greater at higher power and hence higher intensity. The acoustic power (W) represents the intensity emitted by a given surface. For example, with approximately the same [CS2]0 ) 13.3 × 10-4 M and temperature (5 °C) and with air as the irradiating gas, the first-order degradation rate constant, k, of CS2 at 50 W (39.47 W/m2) was about twice that at 14 W (11.04 W/m2), 0.13 vs 0.06 min-1. Similarly, at a temperature of 50 °C, the zero-order degradation rate constant at 30 W (23.68 W/m2) was 50% greater than that at 14 W, 1.8 × 10-5 vs 1.2 × 10-5 M min-1. The power (or acoustic) intensity I, which is proportional to the applied power density, is a function of the acoustic amplitude or pressure, Pa. In the case of a progressive planar or spherical wave, I (in W/m2) is directly related to Pa by eq 3, where F is the

I ) Pa2/2Fc

(3)

density of the fluid (e.g., water) and c is the speed of sound in the fluid (1500 m/s in water), and the term Fc represents the acoustic impedance (Z) of the medium.12 An increase in the ultrasound intensity results in an increase in the acoustic amplitude, which favors more violent cavitation bubble collapse because the bubble collapse time, the transient temperature, and the internal pressure in the cavitation bubble during collapse are all dependent on the acoustic amplitude, Pa. That is, high enough acoustic power results in transient cavitation.20 Hence, the results of an increase in the sound intensity are greater sonochemical effects, resulting in higher CS2 degradation rates. However, Vichare

et al.21 in their theoretical study based on the model of single cavity dynamics using the Rayleigh-Plesset equation discussed the effect of intensity and the frequency of ultrasound on the quantum of energy dissipation and rate of dissipation. They noted that in lowering the intensity, though it resulted in an increase in the estimated energy associated with the individual cavity, the number of cavities could also be substantially less. They also noted that the energy dissipation rate for complete adiabatic collapse conditions decreases initially but increases at higher intensities, although marginally. They concluded that the net effect of this could be an optimal intensity, giving the maximum sonochemical effect. Thus, it could also be concluded that the higher CS2 degradation at higher intensities resulted from an increased number of cavitation events. As illustrated in Figures 2 and 3, the rate of CS2 sonochemical degradation in the presence of the different irradiating gases was in the order He > air > N2O > Ar, at both low temperature (5 °C) and high temperature (50 °C). These results are consistent with our earlier studies10 at 20 °C and are well explained by previous investigators.8-10 Reaction Kinetics As shown in Figures 4 and 5, zero- and first-order plots of CS2 degradation at 5 and 20 °C (air as the irradiating gas) and 5 and 50 °C (N2O as the irradiating gas) both indicate a change in the kinetic order of the reaction with a change in the temperature of the solution. We observed that the reaction rate order is dependent on the temperature (T) range studied: zeroorder at T g 10 °C and first-order at T e 5 °C. The zeroand first-order rate constants for CS2 degradation under the conditions of these experiments are summarized in Table 1. The zero-order plots for the degradation at T ) 10, 15, 20, 30, and 50 °C are shown in Figure 6. The first-order plots for the degradation of CS2 at 5 °C but different initial CS2 concentrations, [CS2]0, and at approximately the same [CS2]0 but different temperatures (i.e., T ) 1, 3, and 5 °C) are shown respectively in Figures 7 and 8. As expected, sonochemical degradation decreases with an increase in the temperature. Hence, the appropriate form of the Arrhenius law11 is applied to the data to obtain the activation energy, EA, for the reactions, and the results are shown in Figures 9 and 10 respectively for the zero- and first-order reaction rate constants. As indicated in Figures 9 and 10, the activation energies for the zero- and first-order degradation of CS2 with air as the irradiating gas were found respectively to be 7.2 kJ/mol at the higher temperatures and 28.7 kJ/mol at the lower temperatures. Because EA represents the potential energy of activation, the higher activation observed for the reactions at the lower temperature range suggests that the higher amount is associated with the cavitation event at the lower temperatures compared with that at the higher temperatures. That is, the factor exp[(bond energy)/RT] that can be achieved at the lower temperature is more significant, and the fraction of the molecules with sufficient energy for dissociation will be higher.16 The kinetics of the sonochemical degradation of pollutants is either first- or zero-order as observed by most investigators. According to the “structured hot spot” model used to explain the results of most studies in environmental sonochemistry, the three regions for

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the occurrence of chemical reactions in a cavitation event are a hot gaseous cavity, an interfacial region surrounding the inner cavity, and the bulk liquid medium at ambient temperature.8,15 The changes in sonochemical activities resulting from temperature changes are believed to be due to the phase in which the reaction leading to chemical degradation occurs and to the physical and chemical properties of the medium (e.g., concentration of the solute).15 The availability and the relative rates of diffusion of free radicals (e.g., •OH) to the reaction zone might then determine the ratelimiting step and the overall order of the reaction. Sehgal and Wang15 in the sonochemical degradation of thymine also observed a change in the reaction order from first to zero as the temperature increased: the reaction rate was zero-order in the temperature range of 10-50 °C and first-order at lower temperatures. They explained the variation of the reaction order with the help of a cavitation-diffusion model proposed by Margulis.22 They noted that the concentration and temperature gradients inside a bubble and in the solution force the free radicals and thymine to diffuse from opposing directions into the bubble-liquid interface, where they react almost instantaneously. As the reaction temperature increased, they suggested that the rate of diffusion of thymine from the bulk liquid to the reaction zone was accelerated. However, the increase in the temperature was simultaneously accompanied by a decrease in the cavitation intensity, reducing the amount of free radicals produced within the bubble. On the other hand, they indicated that, at low solution temperatures, intense cavitation resulted in high intracavity temperatures and high concentrations of radicals, which led to a rapid diffusion of radicals into the interface. Hence, they concluded that the rate-limiting step for the reaction was the diffusion of the substrate (thymine) into the interface at the low temperatures, and at the high temperatures, it was the diffusion of the free radicals into the interface. In general, the results from these studies support the theory that there is an abundance of free radicals, which rapidly diffuse into the interface at low solution temperatures. On the other hand, the diffusion of the substrate (here CS2) in the opposite direction into the interface is slow, and it is, therefore, the rate-determining step. It is, therefore, expected that CS2 would be used up as soon as it diffused into the reaction zone, resulting in a first-order degradation as observed experimentally at the lower temperatures. At the higher solution temperature, cavitation could be inhibited to the extent that the diffusion of free radicals would become the ratedetermining step. The sonochemical degradation reaction is then zero-order with respect to CS2, as illustrated experimentally at the higher temperatures. Reaction Products Analysis of the reaction products indicates that ultrasonic irradiation of CS2 in the presence of air, Ar, and N2O results in its oxidation to mainly sulfate. The formation of sulfate as a function of temperature is illustrated in Figure 11. As expected, the lower the temperatures are, the higher the sulfate production rate due to more severe cavitation events, resulting in the enhanced availability and diffusion of •OH radicals. The formation of sulfate was also enhanced at the lower pH values. With an initial CS2 concentration of (13.6013.85) × 10-4 M and at 20 °C and 50 W, the amount of

sulfate formed after 16 h was about 1.8 × 10-3 M at pH 9 compared with 9.0 × 10-4 M at pH 11 using air as the dissolved gas. For the same conditions but at pH 9 and after 6 h, the amount of sulfate formed was 1.1 × 10-3, 8.3 × 10-4, and 9.1 × 10-4 M respectively in dissolved air, Ar, and N2O. The decrease in the sulfate production rate with an increase in the solution pH can be partly explained by the rapid dissociation of •OH in alkaline solutions as illustrated by eq 4, where kf ) 1.2 kf

•

OH + OH - y\ z O•- + H2O k b

(4)

× 1010 M-1 s-1 for the forward reaction and kb ) 9.3 × 107 s-2 for the backward reaction and the oxide radical ion (O•-) is known to react more slowly with the same substrate than •OH.22 Results of previous sonochemical studies of other reduced sulfur compounds showed similar observations.23,24 Kotronarou et al.23 found the sonochemical oxidation of H2S or S(II-) {[S(-II) ) [H2S] + [HS-] + [S2-]} solutions to proceed rapidly with a zero-order rate, resulting in the formation of sulfate (SO42-) and sulfite (SO32-) as the main products and thiosulfate (S2O32-) as the minor product at pH g 10. They suggested that the apparent zero-order dependence on [S(II-)] was the result of the reaction of the substrate with a •OH radical as the main pathway. They proposed that the rate-determining step in the overall reaction was the reaction of HS- and the oxidation intermediates with a •OH radical in the liquid phase as it diffused out of the cavitation bubble (HS- + •OH f HSOH-). They also observed a decrease in the zeroorder rate constant at pH > 10 and also attributed their observation partly to the dissociation of •OH in the alkaline solutions. In our earlier study we investigated the effects of oxidants, such as H2O2, and •OH radical scavengers, such as 1-butanol, on the product (i.e., sulfate) production rate during ultrasonic irradiation in the presence of air as the irradiating gas.10 We found that the formation of sulfate was enhanced by the addition of H2O2 and inhibited by the addition of 1-butanol, suggesting that the •OH radical played a major role in the sulfate formation mechanism. It is, therefore, conceivable that the rate of formation of an activated complex involving CS2 and •OH determines the overall reaction. Reaction Pathways On the basis of the results discussed, the ratedetermining step in the overall sonochemical oxidation of CS2 in the presence of air (i.e., oxygenated aqueous solutions) to produce sulfate appears to be dependent on the availability of •OH radicals for reactions in the interface of the bubble. The following reaction pathways are proposed:

CS2 + •OH f CS2OH•

(5)

CS2OH• f OCS + HS•

(6)

OCS + H2O f H2S + CO2

(7)

H2S + •OH f HS• + H2O

(8)

HS• + O2 f HSO2•

(9)

HSO2• h SO2•- + H+

(10)

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SO2•- + O2 f SO2 + O2•-

(11)

SO2 + H2O h SO2‚H2O

(12)

SO2‚H2O h H+ + HSO3-

(13)

HSO3- h H+ + SO32-

(14)

2•OH f H2O2

(15)

HSO3- + H2O2 f HSO4- + H2O

(16)

This mechanism indicates that the sonochemical oxidation of CS2 to sulfate proceeds mainly through oxidation by the •OH radical and H2O2 produced from its recombination reactions. In addition, the low EA values ( air > N2O > Ar, at both low and high temperatures. The sonochemical oxidation decreases with an increase in the temperature, and the reaction rate order is temperature dependent: zero-order at T g 10 °C and first-order at T e 5 °C. With air as the irradiating gas, the activation energy, EA, for the zero-order degradation of CS2 was 7.2 kJ/mol at the higher temperatures compared to 28.7 kJ/mol for the first-order degradation at the lower temperatures. The results suggest that the availability and the relative rates of diffusion of free

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radicals (e.g., •OH) to the interfacial reaction zone determine the rate-limiting step and the overall order of the reaction. The product of the reaction in the presence of air (and other dissolved gases, e.g., Ar, N2O) was mainly sulfate, whose formation was enhanced at lower pH values and also at lower temperatures when solutions from experiments in the pH range of 8-11 (T ) 20 °C) and temperatures of 5-20 °C (pH ) 11) were analyzed. Acknowledgment The authors are grateful to Air Force Office of Scientific Research for financial assistance (Grant F49620-95-1-0541) and Department of Energy (Grant DE-FC04-90AL66158) and the equipment support (through Title III) of the Chemical Engineering Department at North Carolina Agricultural and Technical State University. Literature Cited (1) Adewuyi, Y. G. Oxidation of Biogenic Sulfur Compounds in Aqueous Media In Biogenic Sulfur in the Environment. ACS Symp. Ser. 1989, 393, 529. (2) Adewuyi, Y. G.; Carmichael, G. R. 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. (3) Elliot, S. Effect of Hydrogen Peroxide on Alkaline Hydrolysis of Carbon Disulfide. Environ. Sci. Technol. 1990, 24, 264. (4) Hartikainen, T.; Ruuskanen, J.; Martikainen, P. J. Carbon Disulfide and Hydrogen Sulfide Removal with a Peat Biofilter. J. Air Waste Manage. Assoc. 2001, 51, 387. (5) Hugler, W.; Acosta, C. Biological Removal of Carbon Disulfide from Waste Air Streams. Environ. Prog. 1999, 18, 173. (6) Tsai, C.-H.; Lee, W.-J.; Chen, C.-Y.; Liao, W.-T.; Shih, M. Formation of Solid Sulfur by Decomposition of Carbon Disulfide in the Oxygen-Lean Cold Plasma Environment. Ind. Eng. Chem. Res. 2002, 41, 1412. (7) Suslick, K. Sonochemistry. Science 1990, 247, 1439. (8) Adewuyi, Y. G. Sonochemistry: Environmental Science and Engineering Applications. Ind. Eng. Chem. Res. 2001, 40, 4681. (9) Entezari, M. H.; Peeter, K.; Otson, R. The Effect of Frequency on Sonochemical Reactions III: Dissociation of Carbon Disulfide. Ultrason. Sonochem. 1997, 4, 49. (10) Appaw, C.; Adewuyi, Y. G. Destruction of Carbon Disulfide in Aqueous Solutions by Sonochemical Oxidation. J. Hazard. Mater. 2001, B90, 237. (11) Thompson, L. H.; Doraiswamy, L. K. Sonochemistry: Science and Engineering. Ind. Eng. Chem. Res. 1999, 38, 1215. (12) Luche, J.-L. Synthetic Organic Sonochemistry; Plenum Press: New York, 1998. (13) Noltingk, B. E.; Neppiras, E. A. Cavitation Produced by Ultrasonics. Proc. Phys. Soc. London, Ser. B 1950, 63, 674. (14) Neppiras, E. A. Acoust. Cavitation Phys. Rep. 1980, 61, 159. (15) Sehgal, C. M.; Wang, S. Y. Threshold Intensities and Kinetics of Sonoreaction of Thymine in Aqueous Solutions at Low Ultrasonic Intensities. J. Am. Chem. Soc. 1981, 103, 6606. (16) Entezari, M. H.; Kruus, P. Effect of Frequency on Sonochemical Reactions. I: Oxidation of Iodide. Ultrason. Sonochem. 1994, 1, S75. (17) Entezari, M. H.; Kruus, P. Effect of Frequency on the Sonochemical Reactions II: Temperature and Intensity Effects. Ultrason. Sonochem. 1996, 3, 19. (18) Sochard, S.; Wilhelm, A. M.; Delmas, H. Modeling of Free Radicals Production in a Collapsing Gas-Vapor Bubble. Ultrason. Sonochem. 1997, 4, 77. (19) Sochard, S.; Wilhelm, A. M.; Delmas, H. Gas-Vapor Bubble Dynamics and Homogeneous Sonochemistry. Chem. Eng. Sci. 1998, 53, 239. (20) Monnier, H.; Wilhelm, A. M.; Delmas, H. The Influence of Ultrasound on Micromixing in a Semi-batch Reactor. Chem. Eng. Sci. 1999, 54, 2953.

(21) Vichare, N. P.; Senthikumar, P.; Moholkar, V. S.; Gogate, P. R. Energy Analysis in Acoustic Cavitation. Ind. Eng. Chem. 2000, 39, 1480. (22) Margulis, M. Sonochemistry and Cavitation; OPA (Amsterdam) B. V. Gordon and Breach Science Publ.: New York, 1995. (23) Kotronarou, A.; Mills, G.; Hoffmann, M. R. Oxidation of Hydrogen Sulfide in Aqueous Solution. Environ. Sci. Technol. 1992, 26, 2420. (24) Kotronarou, A.; Hoffmann, M. R. The Chemical Effects of Collapsing Cavitation Bubbles: Mathematical Modeling. Adv. Chem Ser. 1995, 244, 233. (25) Sparks, D. L. Environmental Soil Chemistry; Academic Press: New York, 1995. (26) Kang, J.-W.; Hung, H.-M.; Lin, A.; Hoffmann, M. Sonolytic Destruction of Methyl tert-Butyl Ether by Ultrasonic Irradiation: The Role of O3, H2O2, Frequency, and Power Density. Environ. Sci. Technol. 1999, 33, 3199. (27) Serpone, N.; Colarusso, P. Sonochemistry I. Effects of Ultrasound on Heterogeneous Chemical ReactionssA Useful Tool to Generate Radicals and to Examine Reaction Mechanisms. Res. Chem. Intermed. 1994, 20, 635. (28) Riesz, P.; Berdahl, D.; Christman, C. L. Free Radical Generation by Ultrasound in Aqueous and Nonaqueous Solutions. Environ. Health Perspect. 1985, 64, 233. (29) Buxton, G. V.; Greenstock, C. L.; Helman, W. P.; Ross, A. B. Critical Review of Rate Constants for Reactions of Hydrated Electrons, Hydrogen Atoms and Hydroxyl Radicals (•OH/•O-) in Aqueous Solution. J. Phys. Chem. Ref. Data 1988, 17, 513. (30) Margulis, M. A.; Didenko, Yu. T. Energetics and Mechanism of Acoustochemical Reactions. Yields of Hydrogen and Hydrogen Peroxide in Different Aqueous Systems. Russ. J. Phys. Chem. 1984, 58, 848. (31) Hart, E. J.; Henglein, A. Free Radical and Free Atom Reactions in the Sonolysis of Aqueous Iodide and Formate Solutions. J. Phys. Chem. 1985, 89, 4342. (32) Kotronarou, A.; Mills, G.; Hoffmann, M. R. Ultrasonic Irradiation of p-Nitrophenol in Aqueous Solution. J. Phys. Chem. 1991, 95, 3630. (33) Bielski, B. H. J.; Cabelli, D. E.; Arudi, R. L.; Ross, A. B. Reactivity of HO2/O2- Radicals in Aqueous Solution. J. Phys. Chem. Ref. 1985, 14, 1041. (34) Bielski, B. H. J.; Allen, A. O. J. Mechanism of the Disproportionation of Superoxide Radicals. Phys. Chem. 1977, 81, 1048. (35) Bielski, B. H. J. Reevaluation of the Spectral and Kinetic Properties of HO2 and O2- Free Radicals. Photochem. Photobiol. 1978, 28, 645. (36) Liao, C.-H.; Gurol, M. D. Chemical Oxidation by Photolytic Decomposition of Hydrogen Peroxide. Environ. Sci. Technol. 1995, 29, 3007. (37) Farhataziz, R. A. B. Selected Specific Rates of Reactions of Transients from Water in Aqueous Solutions. III. Hydroxyl Radicals and Perhydroxyl and Their Radical Ions; Report No. NSRDS-NBS59; National Bureau of Standards: Washington, DC, 1977. (38) Graedel, T. E.; Wechsler, C. J. Chemistry Within Aqueous Aerosols and Raindrops. Rev. Geophys. 1981, 19, 505. (39) Beckett, M. A.; Hua, I. Elucidation of the 1,4-Dioxane Decomposition Pathway at Discrete Ultrasonic Frequencies. Environ. Sci. Technol. 2000, 34, 3944. (40) Hua, I.; Hochemer, R. H.; Hoffmann, M. R. Sonochemical Degradation of p-Nitrophenol in a Parallel-Plate Near-Field Acoustical Processor. Environ. Sci. Technol. 1995, 29, 2790. (41) Hoffmann, M. R.; Hua, I.; Hochemer, R. Applications of Ultrasonic Irradiation for the Degradation of Chemical Contaminants in Water. Ultrason. Sonochem. 1996, 3, S163.

Received for review January 22, 2002 Revised manuscript received July 24, 2002 Accepted July 29, 2002 IE020069A