Ind. Eng. Chem. Res. 1988,27, 195-197
195
H2S Vapor-Liquid Equilibrium Measurement by the Electrode Method A new method for the measurement of the H2Svapor-liquid equilibrium (VLE) behavior for aqueous solutions using p H and silver sulfide electrodes was developed. The validity of this method was verified by comparing the experimental data on MEA (monoethanolamine)/H2S VLE with literature results. The electrode method showed advantages in the measurement of H2SVLE for very lean solutions, for which the conventional sparging method was not suitable. The VLE data in the very kPa) showed a significant effect due to the reaction of the amine lean loading region (PHs < with water, and these data verified the superiority of the Kent-Eisenberg model over the simple thermodynamic model. Aqueous alkanolamines have been used widely in industry for removing H2S from gas streams (Kohl and Riesenfeld, 1979; Astarita et al., 1983). Knowledge of the vapor-liquid equilibrium (VLE) behavior of H2Sfor these amine systems is important and essential to the design and operation of the gas absorber and stripper. Many investigators have determined H2S VLE for various aqueous amine solutions (Leibush and Shneerson, 1950; Atwood et al., 1957; Muhlbauer and Monaghan, 1957; Jones et al., 1959; Kent and Eisenberg, 1976; La1 et al., 1985). Although these studies covered a wide range of amine concentrations and temperatures, there are very few experimental data in the low loading region encountered in many commercial operations. Recently, the increasing use of high-sulfur-content fuel, the stricter air pollution regulation, and the incentives for sulfur recovery have all demanded removal of H2S to a very low level, which requires a better understanding on the VLE behavior of H2S in the lean loading region. Kent and Eisenberg (1976) developed a model that has correlated previously published data very well and that can predict the VLE for H2S/aqueousamine systems over a wide range of H2S loadings and temperatures. The validity of their model could not be verified for lean loading solutions, however, due to the lack of data in this region. A common method for obtained VLE at low loading is to sparge inert gas into an excess of known solution followed by analysis of the gas for H2S. However, this method had not provided reliable VLE data at very low H2S partial pressures. We have developed a novel method for the measurement of H2S vapor-liquid equilibrium using pH and silver sulfide electrodes, and we found that this method was suitable for H2SVLE in the lean loading region. VLE data for 2.5 M MEA/H2S solution were obtained in a wide range of H2Sloadings. The validity of the electrode method was verified by comparison of the experimental data with literature results. The data in the lean loading region were used to discuss and verify the validity of Kent-Eisenberg's model and the simple thermodynamic model. Theory We determine the H2S VLE by measuring the electropotential difference between a pH electrode and a silver sulfide electrode. The pH electrode responds to the hydrogen ion activity according to the Nernst equation: F In uH+ = constant + -EH+ (1) RT Similarly, the silver sulfide electrode responds to the sulfide activity as follows: 2F In usz- = constant - -EszRT The equilibrium of the H2Sgas with hydrogen and sulfide ions can be expressed by
HZS,,, + 2H+
+ S2-
(3)
Thus, the H2S vapor pressure is given by In PHzs = constant 2 In uHt + In u p
+
(4)
Substitution of eq 1 and 2 into eq 4 gives 2F In PH2S = constant -(EH+ - Esz-) (5) RT This equation shows that the partial pressure of H2S is related to the electropotential difference between the pH electrode and the silver sulfide electrode. At 25 "C, it is
+
log PHzs = constant
+ -----(EH+ 1 0.0296
-
Experimental Section As shown in Figure 1, an Orion Model 901 ion analyzer was used to measure the electropotential difference between a combination Ag/AgCl pH electrode (Orion Model 91-02-00) and a silver sulfide electrode (Orion Model 9416). The electrodes were placed through a rubber stopper into approximately 100 mL of MEA/H2S solution in a thermostated cell. The solution was agitated by a magnetic stirrer, and the temperature was controlled by a water bath. The total sulfide concentration of the MEA/H2S solution was analyzed by silver nitrate titration for each experiment. In the titration, a Metrohm automatic titrator with silver wire and calomel reference electrodes and a 0.05 N AgNO, solution were used. The constant in eq 6 was determined by calibration with a solution of 2.5 M MEA/1.0 M H2S. The vapor pressure of this solution was taken to be 0.655 kPa at 25 "C from the prediction of both the Kent-Eisenberg model and the simple thermodynamic model. Both models were also in good agreement with the available experimental data (Leibush and Shneerson, 1950; Muhlbauer and Monaghan, 1957). The value of the constant in eq 6 was determined on the day of the experiment, and it ranged from -18.0 to -18.6 in our experiments. Rich solutions were prepared by saturating 2.5 M MEA with H2Sat atmospheric pressure. Low loading solutions were prepared by mixing appropriate quantities of the rich solution and an H2S-free 2.5 M MEA solution. To avoid the possible oxidation of H2S,the 2.5 M MEA solution was purged with nitrogen gas for at least 30 min before saturating with H2S or mixing with the rich solution. The electropotential difference measured by the electrodes usually reached steady state in 15-50 min. The values of the measured electropotential differences in our experiments ranged from 430 to 570 mV. Results and Discussion Comparison of Data with Literature Results. A series of experimental runs were performed to measure the VLE behavior for the 2.5 M MEA/H2Ssolutions at 25 "C by using the electrode method. As shown in Figure 2, VLE
0888-5885/88/2627-0195$01.50/0 0 1988 American Chemical Society
196 Ind. Eng. Chem. Res., Vol. 27, No. 1, 1988
I
Reference1
Thermometer
Thermostat&
Simple Thermodynamic Model
0
a
-8
data with good reproducibility over a wide range of H2S loadings (expressed as mol of H2S per mol of MEA) were obtained. The data for H2Spartial pressures between 0.01 and 7 kPa were in good agreement with the available literature data (Leibush and Shneerson, 1950; Muhlbauer and Monaghan, 1957). By using the electrode method, we were able to obtain data at very lean loadings (for H2S partial pressures down to as low as lo4 E a , i.e. 1ppm H2S in the gas mixture of 100-kPa total pressure) where the conventional sparging method had never reached before. The low end of the H2Spartial pressure (about lo4 kPa) for the electrode method was limited by the determination of the H2S loading in the solution (0.003 mol of H2S/mol of MEA) with silver nitrate titration. Comparison of Data with VLE Models. Kent and Eisenberg (1976) developed a model to correlate vaporliquid equilibrium (VLE) data for the absorption of H2S in amines. The model took into account the equilibrium reactions K
(9)
+ S2-
(10)
& H+ + OH-
(11)
HS- _ft3, H+ H20
(7)
& H+ + HS-
H2S,,,
[H2SIO= [H2S] + [HS-] [H+] + [RNH3+] = [OH-]
+ [S2-]
+ [HS-] + 2[S2-]
10-3
10-2
Monaghan's Data
io-'
10
1
Figure 2. Electrode method data that agree with literature results, showing the effect of hydroxyl ion and verifying the validity of the Kent-Eisenberg model in very lean loading regions. Table I. Values of Equilibrium Constants at 25 "C equilibrium const unit value mol/L 1.038 X Kl kPa.L/mol 1.064 X lo3 Hw mol/L 1.052 x 10-7 K2 K3 mol/L 1.818 x 1 0 4 4 Kw mo12/L2 1.067 X
K1, as the adjustable variable and accepted all the remaining equilibrium constants as published. The values of K1 for the amine obtained by their method showed an Arrhenius dependency with temperature. In the present work, we used the Kent-Eisenberg model and accepted the values of all the equilibrium constants used by the authors, as listed in Table I. However, we have found and corrected two typographical errors in their paper: (1)coefficient E for K, is positive, i.e., 0.1361 X 1013 (for T in K), not negative; and (2) coefficient C for K z is negative, i.e., -1.9476 X lo8 (for T i n K), not positive. We also looked into the comparison of the simple thermodynamic VLE model with experimental data. This model considered only the equilibrium reaction H2S,,, + RNH,
and the material and charge balances [RNHZIO = [RNHZ] + [RNHS+]
10-4
(I
y2/(1-y), Where y is H,S Loading (9-mole H2S/g-mole MEA)
Figure 1. Apparatus setup for H2S VLE measurement.
RNH3+2H+ + RNH2
10-5
Muhlbauer This Work
pH@
(14)
The equilibrium partial pressure of H2S was calculated assuming that the vapor pressure of H2S is related to the free H2S concentration in the liquid phase by a Henry's law relationship (McNeil and Danckwerts, 1967): pH@ = HH&[H2Sl (15) Danckwerts (1970) recommended that the Henry's law constant be corrected for highly ionized solutions. However, the method was complicated and required information that was not readily available. Rather than using Danckwerts' approach, Kent and Eisenberg (1976) treated the equilibrium constant representing the amine reaction,
(16)
where the equilibrium constant, Kp, was equivalent to the ratio of the dissociation constants of H2Si1!and the protonated amine (i.e., Kp = K2/K,). The equilibrium partial pressure of HzS was then given by
(12) (13)
2HS- + RNH3+
1 = -HHzS[RNH2]0-
KP
Y2
(1 - Y)
(17)
as can be derived from the stoichiometry and the equilibrium relationahip (Astarita et al., 1983). This model neglects effects of OH- and S2- on the VLE. In Figure 2, equilibrium partial pressures of H2Sfor 2.5 M MEA/HzS mixtures were calculated as a function of the H2Sloading by both models. The modeling results were plotted and compared with literature data (Leibush and Shneerson, 1950; Muhlbauer and Monaghan, 1957) and the electrode method data. As shown in Figure 2, both models fit the experimental data quite well and gave close results for HzS partial pressures down to as low as kPa (i.e., 10 ppm H2Sin the gas mixture of 100-kPa total pressure). For H2S partial pressures below kPa, the simple thermodynamic model predicted lower H2S equilibrium
Ind. Eng. Chem. Res. 1988,27, 197-203
partial pressures than the Kent-Eisenberg model because of a significant effect due to the reaction of the amine with water in the very lean loading region. This significant effect is explained in the following. For the very lean loading region, i.e., for very low partial pressures of H2S or in the absence of H2S, the equilibrium reaction of the amine with water (water acts as a weak acid, and the reaction is the combination of reactions 7 and 11)becomes relatively important: RNH2
+ HzO + RNH3+ + OH-
(18)
This reaction results in a higher OH- concentration for the very lean loading region than for the other region. This reaction makes the total protonated amine (RNH3+)concentration higher than the stoichiometric concentration of the protonated amine from the reaction of the amine with H2S (eq 16). This forces the reversion of reaction 16 and retards the absorption of H2S according to the Le Chatelier principle. Thus, this gives a higher partial pressure of H2S than that predicted by the simple thermodynamic model, which does not take reaction 18 (i.e., the effect of water) into account. However, the KentEisenberg model takes this reaction into account, resulting in good agreement with the data. The electrode method has first provided data to show this effect due to the reaction of the amine with water and verified the superiority of the Kent-Eisenberg model over the simple thermodynamic model in the lean loading region. Conclusions The electrode method for the measurement of H2SVLE was developed and verified. The new method exhibited advantages over the conventional sparging method in H2S VLE measurements, especially for very lean solutions. This method also provided data, for the first time, to show the effect due to the reaction of amine with water, and this verified the superiority of the Kent-Eisenberg model over the simple thermodynamic model in the lean loading region. Nomenclature a = activity E = electrode potential, V
197
F = Faraday’s constant = 96500 C HHzs= Henry’s law constant for H2S, kPa.L/mol K1 = deprotonation constant of amine, mol/L K = stoichiometric equilibrium constant = dissociation constant of H2S, mol/L K 3 = dissociation constant of HS-, mol/L K , = dissociation constant of water, mo12/L2 P = partial pressure, kPa R = gas constant = 8.314 J/K/mol T = temperature, K y = H2S loading, mol of H2S/mol of MEA Subscript
0 = total concentration Registry No. H2S, 7783-06-4.
Literature Cited Astarita, G.; Savage, D. W.; Bisio, A. Gas Treating With Chemical Soluents; Wiley: New York, 1983. Atwood, K.; Arnold, M. R.; Kindrick, R. C. Ind. Eng. Chem. 1957, 49, 1439. Danckwerts, P. V. Gas-Liquid Reactions; McGraw-Hill: New York,
1970.
Jones, J. H.; Froning, H. R.; Claytor, E. E., Jr. J. Chem. Eng. Data 1959, 4, 85.
Kent, R. L.; Eisenberg, B. Hydrocarbon Process. 1976, Feb, 87. Kohl, A. L.; Riesenfeld, F. C. Gas Purification; Gulf Publishing: Houston, 1979. Lal, D.; Otto, F. D.; Mather, A. E. Can. J. Chem. Eng. 1985,63,681. Leibush, A. G.; Shneerson, A. L. J. Appl. Chem. (USSR) 1950,23, 149. McNeil, K. M.; Danckwerta, P. V. Trans. Inst. Chem. Eng. 1967,45, T32. Muhlbauer, H. G.; Monaghan, P. R. Oil Gas J. 1957,55, 139. The University of Texas a t Austin.
* Exxon Research and Engineering Company. Gary T. Rochelle,*’ Philip C. Tsengt W. S. Winston Ho? David W. Savage* Department of Chemical Engineering The University of Texas ut Austin Austin, Texas 78712 and Exxon Research and Engineering Company Annandale, New Jersey 08801 Received for review June 1, 1987 Accepted September 24, 1987
Extinction Phenomena in Countercurrent Packed-Bed Coal Gasifiers: A Simple Model for Gas Production and Char Conversion Rates Global carbon conversion and gas production in a countercurrent moving-bed char gasifier is shown to be determined solely by conditions under which the endothermic carbon gasification reactions are extinguished, in the limit of the large activation energy typical of these reactions. A simplified reaction scheme is employed in which gasification agents C 0 2 and H 2 0 and products CO and H2 are lumped intosingle pseudospecies to develop a simple analytical model for conversion in the gasifier in terms of the asymptotic extinction temperature of the gasification reaction. The assumption of water gas shift equilibrium at this temperature then allows a calculation of detailed product gas composition leaving the reaction zone of the gasifier. The derivation is shown to be valid for chars of sufficiently high reactivity. Model results are compared with those of more comprehensive numerical calculations. Numerical modeling of moving-bed (countercurrent) coal gasifiers such as the Lurgi type has shown that peak temperatures and other derived quantities for the active combustion zone are very sensitive to model assumptions and to parameter values which are generally not wellknown (cf. Arri and Amundson (1978);Yoon et al. (1978)). However, overall conversion rates and effluent gas com0888~5885/88/2627-0197$01.50/0
positions have proven to be remarkably insensitive to model details. These observations led Denn et al. (1979) to formulate a simplified “kinetics-free” model for a onedimensional packed-bed gasifier based on material balances with a specified carbon conversion rate, which employed the assumption of water gas shift reaction equilibrium to calculate the produced gas rate and composition. 0 1988 American Chemical Society
