Environ. Sci. Technol. 2003, 37, 1452-1456
Thiocyanate Wet Oxidation JESU Ä S V I C E N T E A N D M A R I O D IÄ A Z * Departamento de Ingenierı´a Quı´mica y Tecnologı´a del Medio Ambiente, Universidad de Oviedo, E 33071 Oviedo, Spain
Aqueous solutions of thiocyanate were oxidized in a semi-batch reactor at temperatures between 170 and 210 °C and system total pressures ranging from 70 to 100 atm. The initial pH of the solution was set at 12, while initial concentrations have ranged between 20 and 1000 ppm according to the typical concentrations of thiocyanatecontaining wastewater. The kinetic data discussion was based on the establishment of a kinetic regime control that guaranteed the validity of the experimental data and that oxygen excess was assured for all the runs. An attempt to establish a thiocyanate wet oxidation pathway has been made on the basis of the reaction intermediates and final oxidation products identified during the experimental work. SO42- and CO3- have been pointed out as the main final reaction products whereas (SCN)2 and SCNO- are suggested as possible short-life oxidation intermediates. The oxidation reaction results were fitted to a first-order kinetic equation with respect to thiocyanate and to 0.81 with respect to oxygen, with an activation energy of 84.4 kJ mol-1.
Introduction Thiocyanate is one of the major constituents of wastewater from factories for the gasification of coal where various byproducts are formed during the production of gas for fuel, coke, and substances for chemical industries. Composition of such effluents can vary considerably because of differences in process conditions and the type of coal used, with the concentration of thiocyanate typically being between 35 and 1250 ppm (1). Toxic and difficult to eliminate by conventional techniques, thiocyanates are also present in several industrial wastewaters such as metal extraction industries, refineries, or agrochemistry and farm processing plants. Toxicity effects include respiration problems or can even provoke human death because of the formation of toxic gases from contact with acids. Different degradation techniques have been proposed for the treatment of this kind of streams, with the biological treatment being one of the most reliable and most economical option. Thiocyanate biological treatment has been studied for decades (2), but its application range is limited to low pollutant concentrations also giving inhibitory effects in microorganism activity. For this reason, some other techniques, most of them based on oxidation-reduction reactions, have been applied for higher concentrations. In this sense, thiocyanate chemical oxidation, which has received considerable attention in the past as an analytical tool for quantitative determination of thiocyanate, has further * Corresponding author telephone: +34-98-5103439; fax: +3498-5103434; e-mail:
[email protected]. 1452
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become a quite promising alternative for the removal of such pollutant. The oxidation media has been very widely varied (3), but in the past few years hydrogen peroxide (4) has stood out as an effective oxidizing agent. Despite their efectiveness, these proposals are normally expensive and present the additional inconvenience of introducing nondesired substances in the wastewater. So again, the development of new, cheaper, and effective techniques seems to be necessary. Wet oxidation is here presented as an adequate oxidation technique for the treatment of thiocyanate-containing wastewaters. Although it is a relatively recently developed technique in industrial practice, the literature concerning wet oxidation is extensive (5, 6). However, to the best of our knowledge, there have been no studies dealing with thiocyanate wet oxidation specifically. Because of the timely interest and the absence of knowledge in the literature, the aim of the present work is to develop an adequate experimental procedure to demonstrate that thiocyanate can be sucessfully wet-oxidized. In such a way, thiocyanate wet oxidation has been analyzed in a wide range of temperatures, pressures, and initial concentrations, paying special attention to the prevailing regime control, to obtain reliable kinetic data and to make a preliminary analysis of the possible reaction mechanisms.
Experimental Section Apparatus and Procedure. Experiments were completed in a 1-L capacity reactor (Parr T316SS) preceded by a 2-L stainless steel water reservoir (7). To ensure the safety conditions specified when working at high pressure and temperature, the loaded volume in each vessel is about the 70% of the total. The liquid solution was prepared from deionized water and reagent-grade thiocyanate, setting the initial pH of the solution (pH 12) by addition of a predetermined amount of KOH (potassium hydroxide pellets for analysis). The equipment was pressurized and preheated to the desired working conditions while the stirrer speed was adjusted to 8.33 s-1 (500 rpm) for all the experiments. The working pressure was provided by bottled compressed oxygen, with the oxygen flow rate adjusted to 2.33 × 10-5 m3 s-1 and controlled by an electronic mass flow controller (Brooks). First of all, the oxygen was bubbled through the water reservoir in order to become saturated in humidity and then was sparged into the reaction vessel. Once the working pressure, temperature, and oxidant flow rate were attained, a precalculated volume of a thiocyanate-concentrated solution (potassium thiocyanate purissimum) at 20 °C was injected into the reactor by means of the pressure supplied by the bottled compressed oxygen. The injection time was taken as the zero time for the reaction. Liquid samples were periodically withdrawn and analyzed until the concentration was less than 1% of its initial value. The reaction temperature and pressure were maintained during the course of the experiment. Analytical Methods. Thiocyanate concentration was determined using the iron(III) complex method (8) based on the formation of a red complex with iron(III) ions in acidic media whose absorbance has been measured at 480 nm with a Perkin-Elmer spectrophotometer. Cyanide concentration has been measured by means of an Orion 9406 selective electrode using method 413 E included in Standard Methods (8), and sulfate has been analyzed by titrometry (9). The absence of interferences in the determination of thiocyanate concentration caused by reaction products was proved. 10.1021/es020103c CCC: $25.00
2003 American Chemical Society Published on Web 03/06/2003
TABLE 1. Reaction Regime Verification and Relevant Kinetic Data Results run
P (atm) T (°C) CSCNo (ppm) k′ (s-1) CO2,sat (mol/L) k (L0.81 mol-0.81 s-1) 1/kLa (s) 1/kCSCNo (s) Ha
1
2
3
4
5
6
7
100 210 100 3.67 × 10-4 0.088 2.63 × 10-3 0.368 1.39 × 104 3.12 × 10-4
100 200 100 1.83 × 10-4 0.084 1.36 × 10-3 0.385 1.06 × 104 2.23 × 10-4
100 200 1000 4.67 × 10-4 0.084 3.47 × 10-3 0.385 8.57 × 105 9.57 × 10-4
100 200 500 3.50 × 10-4 0.084 2.60 × 10-3 0.385 2.12 × 104 6.93 × 10-4
100 200 200 2.50 × 10-4 0.084 1.86 × 10-3 0.385 5.30 × 104 3.78 × 10-4
100 200 50 1.50 × 10-4 0.084 1.12 × 10-3 0.385 2.12 × 105 1.40 × 10-4
100 200 20 1.33 × 10-4 0.084 9.89 × 10-4 0.385 5.30 × 105 8.27 × 10-5
Results and Discussion The influence of the operating conditions on thiocyanate wet oxidation kinetics has been analyzed in a series of runs varying the working temperature and pressure and testing different initial concentrations, paying special attention to establish a kinetic regime control in all the experiments. First of all, the possibility of removing the thiocyanate present in residual water by using pure oxygen as the oxidant has been tested in a set of preliminary experiments using standard working conditions (200 °C, 100 atm, 100 ppm of initial SCN-, 2.33 × 10-5 m3 s-1 for the oxygen flow, and 8.33 s-1 for the stirring speed with an initial pH of 12). Kinetic Regime Verification. The overall oxidation process can be controlled by two steps: mass transfer of oxygen from the gas phase to the liquid phase and reaction between dissolved oxygen and thiocyanate. Verification for the kinetically controlled regime can be achieved by checking that the initial reaction rate is independent of the speed of the agitation in each experiment or by means of equations R to determine the reaction regime, such as kC SCN - , kLa (for the stoichiometric coefficient equal to 1). The Hatta number, developed from film theory of mass transfer, can also be used as a criteria to establish the regime of control:
Ha )
x
2 β-1 R C SCN kL D kC O 2,sat o (β + 1) O2
/
(1)
No mass transfer limitations existed for the experimental conditions as the kinetic regime is ensured for Ha < 0.02. The diffusivity of oxygen was evaluated using the empirical expression proposed by Wilke and Changssee Reid et al. (10)sand the mass transfer coefficient was correlated according to a expression developed by Calderbank and MooYoung (11). To guarantee the validity of the kinetic data obtained from the experimental runs, kinetic control conditions have to be ensured. Table 1 collects the results of the kinetic regime verification for the runs carried out at the highest working conditions together with the experiments at different initial concentrations. The values for the R and β parameters, which will be later determined, are here employed. The criteria are widely accomplished for the severe working conditions, with the kinetic control being guaranteed with a considerable margin for all of the runs. In all the experiments, the higher values of the Hatta number, corresponding to the beginning of the reaction, were always less than 10-3. As the reaction proceeds, the concentration of thiocyanate decreases leading to a lower value of the Hatta number. Temperature and Pressure Influence. Thiocyanate wet oxidation has been studied in a set of experiments with temperatures ranging from 170 to 210 °C. The rest of the working conditions were fixed as the standard ones, providing the alkaline media in order to avoid unnecessary risks in the case of cyanides formation. The initial thiocyanate concentration of 100 ppm has been chosen because it corresponds
FIGURE 1. Thiocyanate wet oxidation fitting to pseudo-firstorder kinetics at different working temperatures. Symbols denote experimental data: (O) 170, (9) 180, (]) 190, (2) 200, and (/) 210 °C. to the most common value of the average thiocyanatecontaining wastewater concentrations. The pollutant is gradually oxidized from its initial concentration until it is completely eliminated, with the reaction rate depending on the working conditions. For all the runs, even when working at high temperature, a pseudo-first-order kinetic fit is obtained (Figure 1). When the working temperature is 170 °C, the time needed to reach a conversion over 95% is about 900 min being k′ ) 4.83 × 10-5 s-1 (r 2 ) 0.996), whereas only half-time is required to reach the same conversion when working at 180 °C. If the temperature is set at 210 °C, the highest reaction rate is obtained within the selected interval of temperatures (k′ ) 3.67 × 10-4 s-1, r 2 ) 0.993), and a 99% conversion is reached in about 210 min. At the beginning of each experience the pH was set to 12, but during the course of the reaction, it slightly decreased due to acidification of the reaction media as the mechanism later explained shows. However, this variation was of little significance as the pH value was always comprised within 12 and 11.7, even for long reaction times. Using hydrogen peroxide as the oxidant, Christy and Egeberg (3) reported first-order kinetics with a value of k′ ) 1.10 × 10-7 s-1 at 25 °C and pH 7 for a thiocyanate initial concentration of 150 ppm, while the use of peroxomonophosphoric acid (12) as the oxidant at 30 °C gave values around 4.83 × 10-3 s-1 for thiocyanate concentrations between 20 and 150 ppm. These last results are attributed to the acidic media and to the high reactivity of the acid employed. In biological treatment (13), the presence of free radicals caused a relevant increase in the rate of thiocyanate oxidation, but no values for the constant rate were given. However, the values of the kinetic constants obtained here are in the range of values quoted in chemical oxidation processes. In contrast, the rate constants reported in thiocyanate ozonation (14) are higher, 9.00 × 108 M-2 s-1. In a typical case, 500 ppm of thiocyanate dissappeared in 8 min when using a ozone dose of 82 mg/min at pH 7. VOL. 37, NO. 7, 2003 / ENVIRONMENTAL SCIENCE & TECHNOLOGY
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FIGURE 3. Thiocyanate wet oxidation. Fitting to pseudo-first-order kinetics for initial concentrations of SCN- (CSCNo).
FIGURE 2. Thiocyanate wet oxidation at different working pressures. (a) (9) 50, (4) 60, and (b) 70 atm. (b) (]) 80, (2) 90, and (O) 100 atm. Kinetic constants can be assumed to be temperaturedependent following an Arrhenius-type behavior. The equilibrium concentration of oxygen was calculated by means of Henry’s law, taking into account the effect of temperature over the vapor pressure of the water, which modifies the real partial pressure of oxygen of the system from run to run. A value of 84.4 kJ/mol was obtained for the activation energy whereas the pre-exponencial factor was 4.81 × 105 s-1 (r 2 ) 0.991). In the wet oxidation of other pollutants such as p-chlorophenol (15), an activation energy of 91.9 kJ/mol and a pre-exponencial factor of 2.07 × 107 L mol-1 s-1 have been reported. In phenol wet oxidation (7), values of Ea ranging from 45.1 to 112.0 kJ/mol have been proposed, but no comparable data relating to thiocyanate wet oxidation have been found in the literature. When using the experimental apparatus presented here, the energy of activation obtained for thiocyanate wet oxidation is greater than the value obtained for phenol wet oxidation, but the pre-exponential factor has a value of the same order of magnitude. This is in accordance with the fact that thiocyanate oxidation, because of its chemical structure, is slower than the oxidation of other organic pollutants usually present in industrial wastewaters. In the analysis of the pressure influence, the standard working conditions were kept, and experiments at different total pressures from 50 to 100 atm were accomplished. Generally, thiocyanate is gradually oxidized in a single step whose length exclusively depends on the working conditions, but when working at the lower pressures (Figure 2), a short induction period can be distinguished. Increments in the working pressure from run to run reveal the effect of the concentration of oxygen in the liquid phase, leading to higher reaction rates and reducing the reaction time needed to oxidize the thiocyanate. When working at 50 atm, a 95% conversion is reached in approximately 700 min (k′ ) 1.00 × 10-4 s-1, r 2 ) 0.995), but when the total pressure is 100 atm, less than 300 min is required to achieve a similar conversion (k′ ) 1.83 × 10-4 s-1, r 2 ) 0.991). However, working pressure has a less determinant influence on thiocyanate wet oxidation kinetics than working temperature. In the range of pressures considered, the behavior of the system properly fits to a pseudo-first-order mechanism, and even for the experiments carried out at the higher pressures, kinetic regime control has been guaranteed. 1454
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The concentration of oxygen was constant for a given run but varied from run to run as the working pressure and temperature changed. The equilibrium concentration of oxygen was again calculated according to Henry’s law including empirical correlations (16). The oxygen reaction order is determined by correlating the experimental data for oxygen solubility and for the reaction rate constant at different β ), yielding a value of β ) 0.81 working pressures (k′ ) kC O 2 2 (r ) 0.990), with k ) 1.37 × 10-3 (mol/L)-0.81 s-1. In previous works (7), the oxygen reaction order has been shown equal to 1 in phenol wet oxidation, but looking through the literature, values between 0.5 and 1.5 (using Cu2+ as catalyst) have been reported (17-19) although a first order with respect to oxygen has been widely proposed (6, 20) when no catalyst is used. For acetic acid wet oxidation (21), an oxygen reaction order of 1.5 that tended to decrease from 1 to 0 as acetic acid concentration decreased has been indicated, while the wet oxidation of acetonitrile and acrylonitrile (6) lead to values of 0.14 and 0.74, respectively. If inorganic compounds are considered, the range of values is smaller, but the disparity of values for oxygen reaction orders is greater. Thus, first order with respect to oxygen has been reported in nitrile and sulfite (22) oxidation, while the value obtained in cyanides wet oxidation (23) was 0.4. More recently, some authors (24, 25) agreed to propose a reaction order of 1 for oxygen when used to oxidize Fe2+ in aqueous solution, and the same value was found when oxidizing Cu+ with molecular oxygen (26). So, when considering oxidation of inorganic compounds with molecular oxygen, the range of oxygen reaction orders is comprised between 0.4 and 1, with the reaction order for oxygen obtained here (0.81) being smaller than the upper limit, which also is the most common value. Initial Concentration Influence. A set of thiocyanate initial concentrations was selected in a range between 20 and 1000 ppm, using the run carried out at 100 ppm as reference. Standard working conditions were kept for all the runs while the pH was initially set to 12. When thiocyanate initial concentration was 200 ppm, the value for the apparent kinetic constant was k′ ) 2.50 10-4 s-1 (r2 ) 0.998). Further experiments carried out at greater initial thiocyanate concentrations led to increasing kinetic rates k′ ) 3.50 × 10-4 s-1 (r 2 ) 0.996) for 500 ppm and k′ ) 4.67 × 10-4 s-1 (r 2 ) 0.995) for 1000 ppm, reaching a 99% pollutant removal in about 300 min. The degradation takes place in a single stage, which is in accordance with the experience carried out at 100 ppm, and this behavior accurately fits to pseudo-firstorder kinetics (Figure 3). The same behavior has been demonstrated for initial thiocyanate concentrations below 100 ppm. Similar reaction times are required to degrade 1000, 100, or 20 ppm of thiocyanate as the higher the starting pollutant concentration, the greater the reaction rate of the process. This behavior can only be explained by considering that thiocyanate wet oxidation takes place through a reaction
mechanism involving free radicals where the rate of consumption of thiocyanate depends on the initial concentration of thiocyanate (eq 2):
dCSCN0.81 1 ) kf(CSCNo-)C O C SCN- with 2 dt m kf(CSCNo-) ) k′′C SCN - (2) o The fitting of the experimental data values for the initial concentration of thiocyanate from 20 to 1000 ppm yielded m ) 0.37 and k′′ ) 1.51 × 10-2 L0.81 mol-0.81 s-1. According to the mechanism below proposed, the higher the initial concentration of thiocyanate, the higher will be the amount of radicals generated during its oxidation, which possitively contributes to the improvement of the kinetics of the overall process. Pathway for Thiocyanate Wet Oxidation. The mechanistic pathways associated with the evolution of the products of thiocyanate oxidation have been, and still are, an area of strong debate. Among all the different oxidants proposed for thiocyanate chemical oxidation during many years, hydrogen peroxide stands out for the purposes of this research because it has also become an important free radical promoter in several wet oxidation systems (27, 28). In this sense, several authors (4) point at thiocyanogen, (SCN)2, together with hypothiocyanous acid, HOSCN, as the main oxidation intermediates, with HSO4- and HCN being the final products favored by acidic conditions. The initial step of the acidic reaction of H2O2 with SCN- is accounted for by the simple two-electron reduction of H2O2 to H2O and HOSCN. This type of reduction is commonly invoked for peroxide reactions where there is a lack of species present to stabilize “Fentonlike” radicals produced by one-electron reductions. In thiocyanate bioxidation, HCO3- and SO42- have been reported as main final products with no cyanides formation (2), while (SCN)2 and HOSCN appeared as intermediates in SCNoxidation through one-electron-transfer mechanism (13). The alkaline conditions employed in our experimental work may condition one-electron-transfer mechanisms, giving rise to a process where free radicals are involved (29):
have been experimentally proved as a consequence of their extremely short residence time in the reaction bulk:
step 1 SCN- + 3/2O2 + 2OH- f N- + SO42- + H2O (5) In accordance with our qualitative analysis, at least 15 min is required to have a significant amount of sulfate in the bulk liquid, which will remain in the final effluent whereas cyanides quickly disappear. This indicates that thiocyanate does not directly oxidize into sulfate and carbonate, but the oxidation intermediates implicit in step 1 are responsible for the formation of SO42-. For safety reasons, pH needs to be controlled in order to prevent the release of HCN in case of cyanides formation. For instance, Mishra and co-workers (12) detected substantial quantities of cianydes in their works with peroxomonophosporic acid under a pH range from 3.7 to 10.6. Besides precautions and for mechanistic purposes, it is also interesting to know whether cyanides are formed or not under wet oxidation conditions. Thus, when using molecular oxygen as the oxidant and temperatures above 50 °C, very common conditions in wet oxidation processes, if cyanides are formed, they will be readily transformed into cyanogen, CON-, and CO32- (23). In our experimental work, the presence of cyanides has been tested at different reaction times, but no traces of cyanides have been detected at any moment, which can be explained considering the high working temperatures and the great excess of oxygen employed. Although cyanides presence has not been experimentally observed, eq 6 is suggested here as the second step of the proposed mechanism because not only thiocyanate but also cyanides oxidize completely, which eventually produce carbonate, indicated as a final product in reaction 8:
step 2
2CN- + O2 f 2CON-
(6)
SCN- + H2O2 f OSCN- + H2O
(3)
This second step is very quick, and the formed cyanogen reacts with excess oxygen in alkaline media to form carbonate as per step 3, releasing nitrogen to the atmosphere. Step 3 can be considered as the “termination” stage where the final reaction products are formed from the intermediates generated during thiocyanate oxidation:
(SCN)2 + 2OH- T OSCN- + SCN- + H2O
(4)
step 3
According to eqs 3 and 4, hypothiocyanite, OSCN-, and (SCN)2 stand out as the most relevant intermediates. Apparently, the use of two-electron oxidants such as peroxomonophosphoric acid (12) also leads to the formation of thiocyanogen (SCN)2, but in this case, through the reaction between SCN+ species and SCN-. We consider for our experimental system that, if thiocyanogen and hypothiocyanite are formed at any moment, they would readily disappear as a consequence of their instability in aqueous solution because of the severe conditions and the excess of oxygen employed. However, we should not totally discard a contribution of such species to the intrinsic free radical mechanism involved in thiocyanate wet oxidation. In this sense, a three-step general pathway is proposed. First of all, thiocyanate is oxidized by molecular oxygen through a one-electron-transfer mechanism generating thiocyanate free radicals, SCN•, which led to the formation of thiocyanogen and hypothiocyanite as reaction intermediates. Under the prevailing working conditions, these two compounds are very quickly hydrolyzed, yielding CN- and SO42in alkaline media. Their formation and further hydrolysis are implicit in reaction 5, although their presence could not
2CON- + 3/2O2 + 2OH- f 2CO32- + N2 + H2O (7)
If the overall process is considered (eq 8), thiocyanate in alkaline media would react with molecular oxygen to form sulfate, carbonate, and nitrogen. No cyanide presence has been experimentally observed at any time, playing the role of oxidation intermediates through the formation of cyanogen and final products:
2SCN- + 11/2O2 + 6OH- f 2SO42- + 2CO32- + N2 + 3H2O
(8)
Thiocyanate wet oxidation involves radical reactions, and according to the first step of this proposed pathway, the overall reaction rate of the process depends on the amount of radicals generated that are proportional to thiocyanate initial concentration. In connection with the values of the kinetic constants reported in Table 1 and with the prior discussion about initial concentration influence, the higher the initial concentration of such pollutant, the higher will be the concentration of radicals in the bulk liquid and the faster will be the oxidation process. Thus, reaction 5 would be the rate-limiting step because the proliferation of radicals is required to form reaction intermediates, which will lead to VOL. 37, NO. 7, 2003 / ENVIRONMENTAL SCIENCE & TECHNOLOGY
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k′
apparent reaction rate constant
kL
mass transfer coefficient for oxygen in the liquid phase
t
time
n, m, R, β
reaction orders
Literature Cited
FIGURE 4. Wet oxidation of thiocyanate at different initial concentrations. Symbols denote experimental data for 20 (×), 50 (0), 100 (b), 200 (/), 500 (2), and 1000 ppm (]). Solid lines denote model curves according to eq 2. the final products. Any rate enhancement in this step is directly transferred to the rest, achieving a significant enhancement in global kinetics. The experimental curves obtained for thiocyanate wet oxidation at different initial concentrations fitted to the kinetic model proposed in eq 2 are shown in Figure 4, revealing a good agreement for the range of operating conditions tested. All the degradation curves present a similar shape despite departing from different concentration points, which also indicates that the same reaction mechanism prevails independently of the initial concentration of thiocyanate.
Acknowledgments The work upon which this paper is based on was financed by the European Union Contract 7220-EB/004.
Nomenclature
(1) Neufeld, R. D.; Mattson, L.; Lubon, P. J. Environ. Eng. ASCE 1981, 107, 1035-1049. (2) Hung, H.; Spyros, G.; Pavlostathis, D. Water Res. 1997, 31, 27612770. (3) Christy, A. A.; Egeberg, P. K. Talanta 2000, 51, 1049-1058. (4) Figlar, J. N.; Stanbury, D. M. Inorg. Chem. 2000, 39, 5089-5094. (5) Copa, W. M.; Gitchel, W. B. In Standard Handbook of Hazardous Waste Treatment and Disposal; Freeman, H. M., Ed.; McGrawHill: New York, 1989; Section 8.8. (6) Mishra, V. S.; Mahajani, V. V.; Joshi, J. B. Ind. Eng. Chem. Res. 1995, 34, 2-48. (7) Vicente, J.; Rosal, R.; Dı´az, M. Ind. Eng. Chem. Res. 2002, 41, 46-51. (8) APHA, AWWA, WEF. Standard Methods for Examination of Water and Wastewater, 20th ed.; APHA: Washington, DC, 1998. (9) Fritz, J. S.; Freeland, M. Q. Anal. Chem. 1954, 26, 1593-1595. (10) Reid, R. C.; Prausnitz, J. M.; Poling, B. E. The Properties of Gases & Liquids, 4th ed.; McGraw-Hill: New York, 1987. (11) Calderbank, P. H.; Moo-Young, M. B. Chem. Eng. Sci. 1961, 16, 39-54. (12) Mishra, D. K.; Dhas, T. P. A.; Bhatnayar, P.; Gupta, Y. K. Indian J. Chem. 1992, 31, 91-96. (13) Adak, S.; Mazumdar, A.; Banerjee, R. K. J. Biol. Chem. 1997, 272, 11049-11056. (14) Jensen, J.; Tuan, Y.-J. Ozone Sci. Eng. 1993, 15, 343-360. (15) Chang, C. J.; Li, S. S.; Ko, C. M. J. Chem. Biotechnol. 1995, 64, 245-252. (16) Himmelblau, D. M. J. Chem. Eng. Data 1960, 5, 10-15. (17) Willms, R. S.; Balinsky, A. M.; Reible, D. D.; Wetzel, D. M.; Harrison, D. P. Ind. Eng. Chem. Res. 1987, 26, 148-154. (18) Katzer, J. R.; Ficke, H. H.; Sadana, A. J. Water Pollut. Control Fed. 1976, 48, 920-933. (19) Kulkarni, U. S.; Dixit, S. G. Ind. Eng. Chem. Res. 1991, 30, 19161920. (20) Portela, J. R.; Lo´pez, J.; Nebot, E.; Martinez, E. Chem. Eng. J. 1997, 67, 115-121. (21) Levec, J.; Smith, J. M. AIChE J. 1976, 22, 159-168. (22) Wilkinson, P. W.; Doldersum, B.; Cramers, P. H. M. R.; Van Dierendonck, L. L. Chem. Eng. Sci. 1993, 48, 933-941. (23) Mishra, V. S.; Joshi, J. B. Indian J. Technol. 1988, 26, 231-236. (24) Emmenegger, L.; King, D. W.; Sigg, L.; Sulzberg, B. Environ. Sci. Technol. 1998, 32, 2990-2996. (25) Ronnholm, M. R.; Warna, J.; Salmi, T.; Turunen, I.; Luoma, M. Chem. Eng. Sci. 1999, 54, 4223-4232. (26) Papassiopi, N.; Gaunand, A.; Renon, H. Chem. Eng. Sci. 1985, 40, 1527-1532. (27) Kolaczkowski, S. T.; Beltran, F. J.; McLurgh, D. B.; Rivas, J. Process Saf. Environ. Prot. 1997, 75, 257-265. (28) Rivas, F. J.; Kolaczkowski, S. T.; Beltran, F. J.; McLurgh, D. B. J. Chem. Technol. Biotechnol. 1999, 74, 390-398. (29) Walker, J. V.; Butler, A. Inorg. Chim. Acta 1996, 243, 201-206.
a
gas-liquid interfacial area
CO2,sat
oxygen saturation concentration in the liquid phase
Ci
concentration in the reaction media
Ci,o
initial concentration in the reaction mixture
DO2
oxygen diffusivity
Ea
activation energy
Ha
Hatta number
Received for review May 27, 2002. Revised manuscript received September 20, 2002. Accepted January 20, 2003.
k, k′′, kf(SCNo)
reaction rate constants
ES020103C
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