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Kinetics and Mechanism of Oxidation of Hydrogen Sulfide by Concentrated Sulfuric Acid Hui Wang,† Ivo G. Dalla Lana, and Karl T. Chuang* Department of Chemical and Materials Engineering, University of Alberta, Edmonton, Alberta T6G 2G6, Canada
Our previous studies have shown that two reactions are significant when H2S is contacted with concentrated sulfuric acid: H2S+H2SO4fS+SO2+2H2O 2H2S+SO2f3S+2H2O At temperatures from 20 to 60 °C and acid concentrations from 88 to 100 wt %, the rate of the first reaction was measured in terms of the initial rate of pressure drop when H2S and concentrated sulfuric acid were brought together in a volume-constant batch reactor. It is indicated that the reaction is first order with respect to H2S. The global rate constant includes the effects of temperature and acid concentration. The effect of temperature was correlated using the Arrhenius equation. The activation energy and the natural logarithm of the preexponential factor are linearly related to the acid concentration. An illustration of molecular collision at the gas-liquid interface was shown to explain the reaction mechanism. This study also provides a kind of quantitative measurement for the oxidizing ability of concentrated sulfuric acid. Introduction Our previous research has indicated that two consecutive reactions occur when hydrogen sulfide and concentrated sulfuric acid are contacted.1 First, hydrogen sulfide is oxidized by sulfuric acid, forming sulfur, sulfur dioxide, and water.
H2S + H2SO4 ) S + SO2 + 2H2O
DH 2 S
∂2CH2S ∂x2
- rH2S(x,t) )
∂CH2S ∂t
(3)
(1)
Second, hydrogen sulfide reacts with sulfur dioxide, formed in the first step in the acid solution, to produce sulfur and water.
2H2S + SO2 ) 3S + 2H2O
the liquid. In the quiescent liquid, a mass balance over a differential element along the direction perpendicular to the liquid surface gives3
(2)
These two reactions have the potential to become an alternative sulfur removal and recovery process which is flexible, inexpensive, and able to reduce gaseous sulfur emissions to zero. Moreover, it can also remove H2S and water vapor from the gas stream simultaneously. To develop this technology, this paper studies the kinetics of reaction (1), the oxidation of hydrogen sulfide by the concentrated sulfuric acid, using initial rate measurements.2 Theoretical Considerations The reaction between H2S and a sulfuric acid solution is a gas-liquid reaction system. The gas, H2S, reacts with either the liquid, H2SO4, or with SO2, dissolved in * To whom correspondence should be addressed. Telephone: (780)492-4676. Fax: (780)492-2881. E-mail:
[email protected]. † Present address: Department of Chemical Engineering, University of Saskatchewan, Saskatoon, Saskatchewan S7N 5C5, Canada.
Proper experimental design can make the diffusion term become zero so that the accumulation term equals the reaction rate. In the Results and Discussion section, the elimination of diffusion and mass-transfer effects is discussed. When reaction changes the number of moles in the gas phase, it is possible to determine the reaction rate in terms of the rate of pressure change in a constantvolume batch reactor.2 For such a reaction carried out isothermally,
V dPT ) -rH2S δRT dt
(4)
For reaction (1), it would appear that the change in moles in the gas phase is zero because it consumes 1 mol of gas, H2S, and generates 1 mol of gas, SO2. However, when taking initial reaction rate measurements, what we are measuring is the change in total pressure, the sum of the partial pressures of H2S and SO2. Analysis of the electron-transfer direction indicates that SO2 formed from the sulfuric acid molecule is generated in the liquid phase. Our research on the solubility of SO2 in concentrated sulfuric acid4 shows that the volume of the sulfuric acid solution in the reactor is more than capable of dissolving the SO2 produced at the initial moment. Accordingly, any net pressure drop observed at the outset of the reaction
10.1021/ie0200407 CCC: $22.00 © 2002 American Chemical Society Published on Web 11/28/2002
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Figure 1. Schematic diagram of the experimental apparatus: (1) SO2 cylinder; (2) H2S cylinder; (3) regulator; (4) stop valve; (5) SO2 reservoir; (6) H2S reservoir; (7) vacuum pump; (8) pressure gauge; (9) speed-variable motor; (10) reactor; (11) thermocouple; (12) impeller; (13) thermobath; (14) warm water inlet; (15) warm water outlet; (16) magnetic spin bar; (17) magnetic stirrer; (18) pressure transducer; (19) Opto-22; (20) computer.
equals the H2S consumption rate, i.e., the rate of reaction (1). Uncertainties will emerge after the reaction has started. For example, reaction (2) would initiate once SO2 has been produced from reaction (1); the produced sulfur in solid state would block the reaction surface; and the product water would dilute the acid concentration. To eliminate these uncertainties and measure the reaction rate under known conditions, the initial rate method is applied. Based on the above discussion, the reaction rate will be measured by recording the pressure drop against time in a volume-constant batch reactor. The pressure drop versus time curve will be extrapolated to the initial moment, at which point the conditions of the reaction system such as the pressure of hydrogen sulfide, the concentration of the sulfuric acid solution, the area of the clear solution surface are well specified and known, and only reaction (1) occurs. By correlation of these initial rates of reaction, -rH2S0, and initial partial pressure, PH2S0, the order of the reaction and the specified reaction rate can be obtained. Experimental Method The schematic diagram of the experimental apparatus is shown in Figure 1. The sulfuric acid solution of known volume was transferred into the reactor, made of Pyrex glass, and the reactor was connected with the feed system. The air in the system was first evacuated using an Edwards-5 two-stage vacuum pump (Edwards High Vacuum, Oakville, Ontario, Canada), and the solution was heated to the desired temperature. During heating, stirring was applied to keep the temperature homogeneous in the liquid phase. The stirring was stopped before the reaction commenced to maintain a flat interface of known interfacial area. Pure H2S is then
introduced from its reservoir into the reactor to a preset initial pressure. A typical run used 200 mL of the acid solution, and the corresponding volume of the gas phase in the closed volume-constant reactor was 580 mL. After the introduction of H2S into the reactor, the pressure of the gas phase was recorded every 1 s using an Alphaline pressure transmitter (model 1151; Rosemount Inc., Chanhassen, MN) connected with an Opto22 data acquisition system. The resolution of the pressure transmitter was 0.1 psi (0.689 kPa). The pressure data were collected for 1 or 2 min. Then the plot of ∆P, which is the difference between the pressure at any time and the pressure at t ) 0, against time, t, was drawn, and the initial pressure drop rate was obtained in terms of the slope of the curve at t ) 0. The temperatures of the gas and liquid in the closed reactor should be identical and constant. In this study, the temperature of the liquid was controlled by a warm water bath, using a thermomix (model 1460; B. Braun Inc., Toronto, Ontario, Canada). The liquid-phase temperature, read to 0.1 °C, was taken as the reaction temperature. The gas phase was not heated directly and, therefore, was assumed to be at room temperature when the moles of gas were calculated using the equation of state for an ideal gas. In fact, when the temperatures of the gas and liquid differed, the warmer liquid would heat the gas near the liquid, creating a temperature gradient in the gas phase. However, the error thus introduced to the molar number calculation would not be significant because (1) H2S at room temperature was introduced into the reactor just before the reaction started, (2) the gas phase was at room temperature initially and only the gas close to the liquid surface was warmed, and (3) the temperature difference between the gas and liquid never exceeded 40 °C.
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Figure 2. Blank experiment: pressure response when introducing nitrogen into the reactor filled with 200 mL of sulfuric acid of 96.04 wt %. Temperature: 21 °C. Run no.: pd_96_test.
Sulfuric acid solutions at various concentrations were prepared by diluting ∼96 wt % sulfuric acid (Fisher Scientific, Nepean, Ontario, Canada) with distilled water. The 100 wt % acid was prepared by mixing 20% free SO3 fuming sulfuric acid (Acros Organics, Fairlawn, NJ) with a 96 wt % sulfuric acid solution. The concentration of the solutions was determined by titration with a standard 0.1 N sodium hydroxide solution (Fisher Scientific, Nepean, Ontario, Canada), using a 0.1% methyl orange solution (Fisher Scientific, Fair Lawn, NJ) as an indicator. The cylinder of hydrogen sulfide (CP grade) and that of prepurified nitrogen were provided by Praxair Products Inc. (Mississauga, Ontario, Canada). It should be mentioned that the stirrer in the gas phase was designed for studying the kinetics of reaction (2), where H2S and SO2 need to be mixed.5 It was not used in this study. Results and Discussion Blank Experiments. The blank experiments were designed to evaluate the performance of the apparatus. For kinetics studies, what we are mostly concerned with is whether the signals collected and displayed instantly and linearly reflect the pressure in the reactor, because any time delay or signal deformation will introduce errors. Also, what we want to know from the blank experiments is whether there is any factor other than reaction and dissolving gases that contributes to the pressure change. Following the experimental procedure and using nitrogen instead of hydrogen sulfide, blank experiments were carried out. Figure 2 indicates that the apparatus is able to record the pressure change instantly. No time delay or signal deformation and oscillation were observed. The mass balance for nitrogen in both the reactor and reservoir, before and after the introduction, showed no mass loss during the gas introduction. Because N2 is inert to sulfuric acid and also does not dissolve in the solution, we may claim that a pressure change from other than reaction and gas component dissolving is absent. Following the blank experiments, an acid concentration range was selected. Significant interaction between hydrogen sulfide and sulfuric acid only occurs when the concentration of the acid solution is larger than 88 wt %. As shown in Figure 3, the reaction rate in both 87.86 and 30.0 wt % sulfuric acid solutions at room temperature was too slow for this setup to detect; however, that in 96 wt % acid was large enough to measure. Therefore,
Figure 3. Rate of interaction between H2S and sulfuric acid of various concentrations. Temperature: 21 °C. Initial pressure of H2S: 61-65 kPa. Run no.: pd_96_s4, pd_88_t1, pd_30_s1.
the range of sulfuric acid concentration chosen in this study was from 88 to 100 wt %. Effect of Mass Transfer. As mentioned earlier, to measure the reaction rate for a gas-liquid reaction, the influence of mass transfer and diffusion had to be eliminated to render the accumulation rate in the reactor equal to the reaction rate. Because only the initial reaction rate was taken into account, it was important to analyze the mass transfer and diffusion effects at time zero of the measurement. At this moment, the gas phase consisted of pure H2S, ignoring the small amount of vapor from the acid solution, the highest of which in this study was 0.050 kPa, compared to 20 kPa of H2S, the lowest H2S pressure in the runs. Thus, there would be a negligible concentration gradient in the gas phase, and mass transfer by diffusion can be neglected. The liquid phase was at a certain acid concentration, uniform within the liquid. Furthermore, Figure 3 indicates that when the acid concentration decreased as low as 30 wt % and reaction between H2S and a sulfuric acid solution no longer occurred, no significant pressure drop was observed over a reasonably long period of time, implying that no processes other than reaction would contribute to the pressure drop. Even though the solubility of H2S in the diluted sulfuric acid was studied by a number of researchers,6,7 the dissolving rate in this study can be ignored. Based on this analysis, at time zero, mass transfer and diffusion in the liquid phase were also insignificant. Therefore, the pressure drop that was measured only originated from reaction (1). Effect of the Reaction Area and Volume. For a gas-liquid reaction, the gaseous reactant must contact the liquid phase to facilitate reaction. Because the interface is where the reactant molecules in the different phases meet and react, its area should influence the reaction rate. To determine how the interfacial area and the volume of liquid-phase affect the reaction rate measurement, two reactors of different shapes were used: a cylindrical reactor which changes the acid solution volumes but keeps the interface area constant and a conical reactor in which different surface areas of the acid solution result from changing different acid volumes. First, the results obtained using the cylindrical reactor, as shown in the upper three rows in Table 1, indicate that the reaction rates at unit H2S pressure (Pa) were nearly identical no matter how the volume of the sulfuric acid solution changed. Second, in the conical reactor, the reaction rate was proportional to the interfacial area, shown in the lower four rows of Table
Ind. Eng. Chem. Res., Vol. 41, No. 26, 2002 6659 Table 1. Values of kP1 for Different Surface Areas and Volumes of the Sulfuric Acid Solutiona
run no.
surface area, m2
acid volume, mL
rate × 1010, mol s-1 Pa-1
kP1 × 108, mol s-1 Pa-1 m-2
pd_96_a1 pd_96_a2 pd_96_a3 pd_96_a4 pd_96_a5 pd_96_a6
0.004 42 0.004 42 0.004 42 0.005 01 0.008 64 0.0119
200 300 400 800 500 300
2.83 2.86 2.81 3.22 5.33 7.63
6.40 6.48 6.35 6.42 6.30 6.41
a
Acid concentration: 96.04 wt %. Reaction temperature: 21.5
°C.
Figure 5. Arrhenius plots of the reactions at different acid concentrations: (b) 99.97 wt %; (9) 96.04 wt %; (2) 92.97 wt %; (1) 91.53 wt %; (O) 90.17 wt %; (0) 87.68 wt %. Table 2. Values of the Preexponential Factor and Activation Energy for Various Acid Concentrations
Figure 4. Plots of the initial H2S consumption rate vs H2S initial pressure. Acid concentration: 96.04 wt %. Run no.: pd_96_t1, pd_96_t2, pd_96_t3, pd_96_t4.
1. Identical rate constants per unit area were obtained from all of the runs. The reaction rate is proportional to the interfacial area but independent of the acid volume, indicating that reaction (1) likely occurs only at the surface of the liquid phase. Order of the Reaction with Respect to H2S and Sulfuric Acid. At a fixed temperature and acid concentration, the initial reaction rate was measured at various initial H2S pressures, following the experimental procedure described. The order of the reaction with respect to the pressure of H2S was obtained by correlating the initial rate and initial pressure. The result for a 96.04 wt % sulfuric acid solution, as an example, is shown in Figure 4, clearly illustrating first-order behavior with respect to H2S at various temperatures. Experimental results from other acid concentrations were shown in ref 8. Analysis of the data indicates that, for all concentrations from 88 to 100 wt % and temperatures from 20 to 60 °C, first-order behavior with respect to the H2S pressure was observed. The slope of the lines of the reaction rate per unit interface versus H2S pressure, obtained by regression provided the rate constant or specific reaction rate. It should be realized that this rate constant was an apparent rate constant, in which the influence of the acid concentration as well as that of the temperature on the reaction rate was involved. This rate equation takes the form of eq 5.
rH2S1 ) kP1aPH2S
(5)
At a given H2S pressure, e.g., 35 kPa, and temperature, e.g., 25 °C, the reaction rate per unit interfacial area for various acid concentrations was calculated using eq 5. The poor correlation obtained between the reaction rate and the molarity of H2SO4 does not suggest a reaction order with respect to sulfuric acid.
acid concn, wt %
preexponential factor, ln A0
activation energy, kJ mol-1
regression coefficient, r2
99.97 96.04 92.97 91.53 90.13 87.68
2.89 ( 0.17 -3.12 ( 0.002 -5.30 ( 0.45 -7.93 ( 0.89 -9.56 ( 0.17 -11.5 ( 0.38
45.1 ( 0.42 32.9 ( 0.004 30.0 ( 1.16 24.1 ( 2.34 21.8 ( 0.45 17.3 ( 1.03
0.9999 0.9999 0.9956 0.9725 0.9906 0.9929
Effect of the Temperature. The Arrhenius plots of the global rate constants, kP1, for different acid concentrations are shown in Figure 5. Correlating the natural logarithm of kP1 and the reciprocal of absolute temperature, T, yields the values of the preexponential factor, A0, and the activation energy, Ea, of the reaction in the acid solution of different concentrations. The values of A0 and Ea for each acid concentration are listed in Table 2. The relationship between kP1 and temperature T could be described by
kP1 ) A0 exp(-Ea/RT)
(6)
Effect of the Acid Concentration on the Reaction Rate. Although the order of the reaction with respect to sulfuric acid could not be expressed with a simple integer, to apply the reaction rate equation in the reactor design, the effect of the acid concentration on the reaction rate should be described quantitatively. As shown in eq 5, the effect of acid concentration, Ca, is involved in the global rate constant, kP1, which, in turn, can be decomposed to the preexponential factor, A0, and activation energy, Ea. The regression analysis between ln A0 and Ca, and between Ea and Ca, respectively, gave rise to the following linear relationships:
ln A0 ) a + bCa
(7)
Ea ) c + dCa
(8)
Figure 6 shows the regression results, and Table 3 lists the values of the parameters a, b, c, and d, and the value of the regression coefficients as well, showing a good correlation. This correlation between the rate constant and the acid concentration in weight percentage is empirical, because the overall concentration does not represent the real amount of species existing in the solution. Equations 7 and 8 indicate that both ln A0 and
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Figure 6. Correlation between ln A0 and Ca and between Ea and Ca.
Figure 7. Rate constant, kP1, and molarity of the species in sulfuric acid solutions. (The molarity curves were reproduced from ref 8.) Table 3. Values of a, b, c, d, and r2 in Equations 7 and 8 eq 7 for ln A0 eq 8 for Ea
intercept
slope
r2
a ) -116 ( 6.8 c ) -178 ( 17
b ) 1.18 ( 0.07 d ) 2.23( 0.18
0.9847 0.9751
Figure 8. Plots of the reaction rate vs the activity coefficient of H2SO4. The reaction rate at 25 °C and 35 kPa of PH2S was calculated using eq 5, and the activity of H2SO4 was found in ref 10.
the ions in the sulfuric acid solution. Then, correlating kP1 with some physical and chemical properties of sulfuric acid such as the electrical conductivity,10 the acidity function,11 and the activity coefficient11 shows that only ln kP1 and the activity coefficient of sulfuric acid are in good correlation (Figure 8). Although this good linear relationship between them cannot be explained from the kinetics or mechanistic point of view, it at least supports the view that molecular H2SO4 present in the acid solution is the active species of the oxidation of hydrogen sulfide. When hydrogen sulfide is oxidized, its sulfur atoms increase their oxidation number by transferring their electrons to the molecules of the oxidizing agent.11 When the gas, H2S, contacts the concentrated sulfuric acid, it attacks the H2SO4 molecule and forms an active intermediate, through which the electrons are transferred and the products, sulfur, sulfur dioxide, and water, are formed. This interaction is illustrated by the consecutive steps (11)
Ea decrease as the acid solution becomes more dilute. As we know, a decrease in the activation energy may lead to an increase in the reaction rate, but overall the value of kP1 decreases as the acid concentration was reduced. Combining eqs 6-8,
kP1 ) ea+bCa-(c+dCa)/RT
(9)
Rearrangement of eq 9 leads to
kP1 ) e[a-(c/RT)]+[b-(d/RT)]Ca ) eR+βCa
(10)
where R and β are constants at a constant T. These forms of the expression of kP1 show the effects of T and Ca clearly and are easily used in the reaction rate calculation. Active Species in Sulfuric Acid to H2S Oxidation. A thermodynamic analysis has suggested that only the reaction between H2S and molecular H2SO4 is likely.1 Plotting both ln kP1 and concentrations of individual species present in the sulfuric acid solution against the weight percentage of H2SO4, as shown in Figure 7, indicates that the trend behavior of ln kP1 is similar to that of the molarity9 of molecular H2SO4, excluding the possibility of reactions between H2S and
Durrant and Durrant13 pointed out that, for the molecule H2S, the distance, H-S, is 1.34 Å and the angle of HSH is 92.2° and, for H2SO4, the length of the πd bonds between O and S is 1.40 Å and the angle between the two πd bonds is 124°. These structure parameters indicate that the attack in step 1 in eq 11 is reasonable and stereofavorable. After an effective collision, an intermediate forms, in which the electrons are transferred directly from the sulfur atom in H2S to the sulfur atom in H2SO4 and the products are formed. Compared
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to steps 2 and 3, step 1 is assumed to be the ratecontrolling step, because more energy is needed to have a reactive collision, forming the active intermediate. Equation 11 is not a mechanistic expression. It agrees with the first-order behavior of H2S, but it is not able to explain the behavior of the sulfuric acid concentration, as mentioned previously. The oxidizing ability of sulfuric acid depends strongly on its concentration. As a result, sulfuric acid at different concentrations may behave like different reactants, leading to reaction rates that do not correspond to a simple order with respect to sulfuric acid. Measurement of the Oxidizing Ability of Sulfuric Acid. Because the sulfuric acid solution plays the role of an oxidizing agent in its reaction with H2S, the values of kP1 from this study can be taken as a kind of quantitative measurement of the oxidizing ability of sulfuric acid in reaction (1), a topic rarely discussed quantitatively either in textbooks13 or in the literature.10,14 Derived from eq 10, the ratio of the oxidizing ability of sulfuric acid at two different concentrations, Ca1 and Ca2, can be described quantitatively as
kP1,Ca1 kP1,Ca2
)e
β(Ca1-Ca2)
(12)
and the comparison of the oxidizing ability can be extended to between two temperatures using eq 10, resulting in eq 13.
kP1,T1 kP1,T2
) e[(c+dCa)/R](1/T2+1/T1)
(13)
For instance, from eq 12, one can estimate that at 25 °C the 95 wt % sulfuric acid solution is 4 times as strong as the 90 wt % sulfuric acid solution in oxidizing ability, and from eq 13, one shows that the oxidizing ability of the 90 wt % sulfuric acid solution at 125 °C is 8.4 times as strong as that at 25 °C. Conclusions The oxidation of hydrogen sulfide by concentrated sulfuric acid behaves as first order with respect to H2S. Sulfuric acid plays the role of an oxidizing agent, which affects the reaction rate through a change of the preexponential factor or the collision frequency, A0, and the activation energy, Ea, of the reaction when its concentration changes. The correlation equations between ln A0 and the acid concentration and between Ea and the acid concentration have been obtained, and both show linear relationships. The reaction likely occurs between H2S and molecular sulfuric acid. The dependence of the rate constants on temperature fits the Arrhenius equation. The reaction probably initiates with gas-liquid molecular collisions at the interface. The value of the rate constant may be used to compare the oxidizing ability of sulfuric acid in reaction (1) at different concentrations and various temperatures. Acknowledgment The authors acknowledge financial support from the Natural Sciences and Engineering Research Council of Canada and beneficial discussions with Dr. Qinglin Zhang. Nomenclature A0 ) preexponential factor, mol s-1 m-2 Pa-1 a ) surface area (eq 5), m2
a ) regression constant arisen from eq 7 b ) regression constant arisen from eq 7 C ) concentration of the reactant, mol m-3 Ca ) apparent sulfuric acid concentration, wt % c ) regression constant arisen from eq 8 d ) regression constant arisen from eq 8 D ) diffusivity, m2 s-1 Ea ) activation energy, kJ mol-1 kC ) specific reaction rate in terms of the molarity concentration, s-1 kP ) specific reaction rate in terms of the pressure of the reactants, mol s-1 m-2 Pa-1 N ) number of moles P ) pressure, Pa, kPa R ) gas constant, J mol-1 K-1 r ) reaction rate (eq 3), mol s-1 m-3 r ) reaction rate (eqs 4 and 5), mol s-1 T ) temperature, K V ) volume of the gas phase, m3 t ) time, s x ) direction perpendicular to the gas-liquid interface, m Subscripts a ) acid T ) total 0 ) Refer to time zero 1 ) Refer to reaction (1) 2 ) Refer to reaction (2) Greek Symbols R ) constant arisen from eq 10 β ) constant arisen from eq 10 δ ) The change in the number of moles of gas reactants per mole of H2S in the reaction
Literature Cited (1) Zhang, Q.; Dalla Lana, I. G.; Chuang, K. T.; Wang, H. Reactions between Hydrogen Sulfide and Sulfuric Acid: A Novel Process for Sulfur Removal and Recovery. Ind. Eng. Chem. Res. 2000, 39, 2505. (2) Fogler, H. S. Elements of Chemical Reaction Engineering; Prentice Hall: Englewood Cliffs, NJ, 1986. (3) Danckwerts, P. V. Gas-Liquid Reactions; McGraw-Hill Book Company: New York, 1970. (4) Zhang, Q.; Wang, H.; Dalla Lana, I. G.; Chuang, K. T. Solubility of Sulfur Dioxide in Sulfuric Acid of High Concentration. Ind. Eng. Chem. Res. 1998, 37, 1167. (5) Wang, H.; Dalla Lana, I. G.; Chuang, K. T. Kinetics of Reaction between Hydrogen Sulfide and Sulfur Dioxide in Sulfuric Acid Solutions. Ind. Eng. Chem. Res. 2002, 41, 4707. (6) Alexandrova, M. V.; Yaroshchuk, E. G. Solubility and GasLiquid Equilibrium in the Systems of Nitrogen-Carbon Disulfide and Hydrogen Sulfide-Acidic Salt Solutions at 50 °C. J. Appl. Chem. 1978, 51, 1221. (7) Douabul, A. A.; Riley, J. P. The Solubility of Gases in Distilled Water and SeawatersV. Hydrogen Sulfide. Deep-Sea Res. 1979, 26A, 259. (8) Wang, H. A Study of the Gas-Liquid Reaction System of Hydrogen Sulfide and Sulfuric Acid. Ph.D. Thesis, University of Alberta, Edmonton, Alberta, Canada, 2002. (9) Young, T. F.; Walrafen, G. E. Raman Spectra of Concentrated Aqueous Solutions of Sulfuric Acid. Trans. Faraday Soc. 1961, 57, 34. (10) Sander, U. H. F.; Fischer, H.; Rothe, U.; Kola, R. Sulfur, Sulfur Dioxide and Sulfuric AcidsAn Introduction to their Industrial Chemistry and Technology; More, A. I., English Ed.; The British Sulfur Corp. Ltd.: London, 1984. (11) Liler, M. Reaction Mechanisms in Sulfuric Acid and Other Strong Acid Solutions; Academic Press: New York, 1971. (12) Weil, E. D.; Randler, S. R. Sulfur Compounds. KirkOthmer Encyclopedia of Chemical Engineering, 4th ed.; John Wiley & Sons: New York, 1997.
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(13) Durrant, P. J.; Durrant, B. Introduction to Advanced Inorganic Chemistry, 2nd ed.; John Wiley & Sons: New York, 1970. (14) Muller, T. L. Sulfuric Acid. Kirk-Othmer Encyclopedia of Chemical Engineering, 4th ed.; John Wiley & Sons: New York, 1997.
Received for review January 16, 2002 Revised manuscript received October 4, 2002 Accepted October 4, 2002 IE0200407
