Reduction of Sulfur Dioxide with Methane over Selected Transition

David J. Mulligan and Dimitrios Berk*. Department of Chemical Engineering, McGill University, Montreal, Quebec, Canada H3A 2A7. The reduction of sulfu...
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I n d . E n g . C h e m . Res. 1989, 28, 926-931


Reduction of Sulfur Dioxide with Methane over Selected Transition Metal Sulfides David J. Mulligan and Dimitrios Berk* Department of Chemical Engineering, McGill University, Montreal, Quebec, Canada H3A 2A7

T h e reduction of sulfur dioxide to produce elemental sulfur is an alternative process for treating SO2 produced in smelters. This process can be made more economical by using a more selective catalyst for the production of elemental sulfur than the traditionally used alumina catalyst. T h e reduction of SOz and CH4 over the transition metal sulfides FeS, MoSz, and WS2 was studied in a gradientless, tubular packed-bed reactor. Experiments were carried out in the temperature range 650-750 "C. T h e inlet molar feed ratio of SOz to CHI was varied between 0.5 and 2.0. FeS showed sulfur selectivities ranging from 8.0 t o 180.0 but formed nonstoichiometric compounds. MoSz also provided similar selectivities and rates of SOz consumption equivalent to activated alumina. T h e selectivities with WS2 ranged from 7 t o 20, and SO2 consumption rates were twice those for alumina or MoSz. However, byproduct production, including the production of elemental carbon, was significantly higher for WS2 in comparison with MoS2. From the results, it was concluded that MoS2 was the most effective of the catalysts considered. As evidence demonstrating the detrimental effects of sulfur dioxide emissions on the environment is accumulated, the necessity of its economic disposal becomes more obvious. A t the present time, most of the SO2 that is recovered from the waste gas streams from the roasting of sulfide mineral ores is oxidized to sulfur trioxide to produce sulfuric acid. An alternative process which has been implemented by the industry in the past is the reduction of SO, to elemental sulfur. As an end product, sulfur has several advantages over sulfuric acid because it can be handled, stored, and transported safely, with no danger to the environment. Sulfur dioxide can be reduced by a number of reducing agents including hydrocarbons, carbon monoxide, hydrogen, and carbon (Rosenberg et al., 1975). However, because of its availability and relatively low price, methane has been the reducing agent in all large-scale industrial processes. The primary reaction between SO, and CH, is 2S02

+ CH,



+ 2[S] + COS


where [SIrepresents the various sulfur species in the gas phase. In addition to the primary products, a number of side reactions may result in the production of undesired sulfur byproducts, H,O, COS, and CS2. Other possible byproducts of the above reaction system include CO, H,, and elemental carbon. Reduction processes described in various papers and patents involve a two-stage process (Hunter and Wright, 1972; McMillan, 1971). In the first stage, methane reduces sulfur dioxide to elemental sulfur and the undesired sulfur byproducts at a temperature greater than 800 "C. After the removal of sulfur, the byproducts are then sent to the second stage, a Claw conversion unit, where the remaining SO2 reacts with H2S and COS to recover additional elemental sulfur. In order to reduce the capital equipment and operating costs of the plant, it is desirable to reduce the size of the Claus conversion stage, or eliminate it entirely, by reducing the amount of byproducts produced in the first stage. The kinetics of the reaction between SO2 and CH, in the presence of an alumina catalyst were studied by Helmstrom and Atwood (1978). Two rate expressions based on a single-site model and a double-site model were reported. Also it was found that the formation of undesired byproducts H2S, COS, and CS, could be minimized by using temperatures ranging from .550 t o 650 "C and an inlet 0888-5885/89/2628-0926$01.50/0

SO,-to-CH, ratio above 2.0. However, under such reaction conditions, the reaction rate is low. For an industrial process for the reduction of SO2 using alumina catalyst, reaction temperatures would have to be higher to increase the reaction rates. Sarlis and Berk (1988) reported rates of production of elemental sulfur and the other reaction products at temperatures between 650 and 750 "C at molar feed ratios between 0.5 and 2 and found low selectivities. Therefore, it is desirable to limit the production of byproducts by using a catalyst that has a higher selectivity than alumina without reducing the primary reaction rate. The purpose of the work that is presented in this paper is to evaluate catalysts for the reduction of SO2with CH4 and compare them with the commonly used alumina catalyst. The evaluation and comparison is based on the reaction rates and selectivity toward the production of elemental sulfur. The catalysts selected for study include MoS2, WS2, and iron sulfide. All three of these sulfides are known to catalyze the decomposition of H2S (Chivers et al., 1980). MoS, is also currently being used as a hydrodesulfurization catalyst (Joffre et al., 1986). MoSz and WS2 do not change sulfidation state; however, sulfur may dissolve in the FeS crystal, forming a nonstoichiometric iron sulfide (pyrrhotite). The activity and selectivity of' each catalyst was measured as a function of temperature and inlet molar feed ratio of SO2 to CH,.

Materials and Methods Catalyst Preparation. MoS2 and WS2 were purchased as 300-mesh powders from Johnson Matthey, Inc., Melvern, PA. The powders were formed by compression into pellets at 4500 psi in a stainless steel die. The pellets were in the shape of disks, 6 mm in diameter and an average 3 mm in thickness. The average weight for the MoS, pellets was 0.37 g and for the WS2 pellets was 0.49 g. Before use in experiments, the pellets were conditioned by placing them in a flow of argon in the reactor tube at a temperature of 700 "C for a period of 6 h. The iron sulfide was prepared from crystalline pyrite (FeS2),which was purchased from Hawthorneden. The crystal was crushed into 8-12-mesh granules, which were desulfurized in the reaction tube at 750 O C in the presence of flowing argon until sulfur ceased to evolve. The granules were then pulverized, and the resulting powder was pelletized at 4500 psi. In this case again, the pellets had the shape of disks and the same thickness as the MoS, and D 1989 American Chemical Societ)

Ind. Eng. Chem. Res., Vol. 28, No. 7 , 1989 927 Table I. Composition of the Inlet Gas Mixtures at Different Feed Ratios SOz/CH4 feed ratio soz, % CH,, % 0.5 20 40 1.0 25 25 1.5 36 24 2.0 30 15

Table 11. Surface Area Analysis surface area initial, m2/g final, m2/g MoSZ 6.68 3.68 3.00 2.25 wsz FeS 1.2-3.6 2.20-2.29

WS2 pellets. However, the die that was used for the pelletization of FeS had a diameter of 12.5 mm because the gritty nature of the FeS powder caused the smaller die to jam. Because of this large size, the iron sulfide pellets had to be quartered in order to obtain an external surface area comparable to that of the MoS2 and the WS2 pellets. Experimental System. Sulfur dioxide, methane, and argon, the carrier gas, were delivered a t a pressure of 20 psig. The flow rates of each of the gases were measured by calibrated rotameters and controlled by needle valves. The gases were then mixed and delivered to the reactor in stainless steel tubing. The reactor was a 65-cm-long, 2.5-cm-i.d. quartz tube. It was heated in a single-zone, heavy-duty, Lindberg Model 1500, tubular furnace. The catalyst was located in the middle of the tube where the temperature was kept uniform. The remainder of the tube was filled with quartz chips to improve mixing and to reduce the void volume. The temperature of the reactor bed was measured by three thermocouples (chromel-alume1 Type K). To verify the uniformity of temperature along the catalyst bed, the thermocouples, which were inserted through the entrance of the reactor along the centerline, were positioned at the beginning, in the middle, and at the end of the catalyst bed. As the gases left the exit of the reactor, they flowed through a U tube, cooled by an ice bath. In this trap, sulfur, water, and any CS2that may have been formed were condensed. The flow rate of the remaining gases was then measured by a bubble flowmeter at ambient conditions. The gas samples were taken directly from the reactor exit line by gas-tight syringes and injected immediately into the gas chromatograph. The gas chromatograph (HP Model 5780) was fitted with two columns. The first one was a Poropak QS column, 45 cm long. The second was a molecular sieve 5A column of the same length. A thermal conductivity detector was used to detect the gases. The detector signal was processed by a H P 3390 integrator. With this system, COz, H2S, COS, SOz, Ar, CH,, and CO could be measured quantitatively, while sulfur, H 2 0 , C02, and elemental carbon could be determined by elemental balance. In addition, the presence of H2 and CS2 could also be detected. Experimental Procedure. Before and after experimentation, the bulk composition of each of the catalysts was determined by X-ray diffraction. In addition, the surface area was measured by the BET method. This was accomplished by using a Micromeritics Flowsorb Model 5200 surface area analyzer. Experiments were carried out in the temperature range from 650 to 750 "C at 25 "C intervals. The ratio of inlet S02-to-CH4concentrations was varied between 0.5 and 2.0. Table I shows the concentrations of each of the gases for the specific ratio used. At the beginning of each experiment, the gas flow rates were set by the rotameters. The inlet concentrations were verified by GC analysis. The volumetric flow of the gases was kept between 5 and 7 cm3/s (at 25 "C and 1 atm), which was above the experimentally determined threshold for external film diffusion. The reaction mixture was sent through the reactor as the temperature stabilized to the

Table 111. Selectivities of Iron Sulfide feed ratio 650 "C 675 "C 700 "C 0.5 46.0 38.0 35.5 48.0 44.0 1.5 180.0

% change

-45 -24

725 "C 9.0 18.0

750 "C 8.0 12.0

reaction temperature. At intervals of 15 min, samples were taken from the exit stream and analyzed until three consecutive analyses yielded results within 5%. At that point, steady state was assumed to be achieved. This procedure was then repeated for each set of reaction conditions, for each catalyst. The reaction rates of each species were calculated by mass balances as described by Mulligan (1988).

Results and Discussion Preliminary experiments showed that the only reaction products were C02,H2S, H 2 0 , COS, CO, sulfur, and elemental carbon. No H2 or CS2 was detected. Also, when no catalyst packing was present in the reactor, no conversion was detected at any of the experimental conditions, confirming the absence of homogenous reaction between reactants. In addition, product stream samples were taken at the exit of the catalyst bed and at the exit of the reactor tube. Comparison of the two samples showed no difference in composition, confirming the absence of homogeneous reactions among product stream components. Finally, replicate experiments showed that conversion, selectivity, and reaction rate were independent of catalyst packing or catalyst samples from batches produced at different times. The surface areas of the catalysts, MoS2 and WSz, decreased during the initial experimentation period. MoS2 lost approximately 45% of its original surface area, while WSz lost only 24%. Both surface areas stabilized at this point (Table 11). The surface area of iron sulfide showed a variation due to problems in degassing the sample before analysis. During this procedure, which was ordinarily carried out at a temperature of 200 "C, sulfur evolved. As the degassing temperature was increased, the quantity of sulfur evolving increased as did the surface area. Since a stable surface area measurement could not be obtained for FeS, the reaction rates based on the surface area for this catalyst were not calculated. However, sulfur selectivities and carbon dioxide yield, which are independent of the surface area, are included. Sulfur-Containing Species. For the purpose of comparing catalysts, selectivity (s) has been defined as follows: r(S) s = r(H,S) + r(C0S) where r(i) is the rate of production of species i in mol/ (mz.s). In order to take into account the difference in surface areas for the different catalysts, the reaction rate data are reported per unit surface area. Table 111 and Figures 1and 2 show the selectivities of iron sulfide, MoSz, and WS2, respectively, at various temperatures and inlet ratios of SO2 to CH,. The selectivities for iron sulfide are significantly higher than those found for WSz or MoSz at temperatures below 725 "C. However, XRD analysis of the iron sulfide before reaction showed that the primary component was FeS, triolite. After reaction, the primary component was found

928 Ind. Eng. Chem. Res., Vol. 28, No. 7 , 1989


~ ~ - _ _




IO i'olir

.5 k e d ?ut o ~














-ernFerature ' 3

Figure 1. Effect of molar feed ratio on the selectivity of MoS, at different temperatures

Figure 3. Effect of temperature on the selectivity of WS, at different molar feed ratios









I b.














p i



I ' o l z r 'eed







Figure 2. Effect of molar feed ratio on the selectivity of WS, a t different temperatures.

to be FeS,,l. This indicates that iron sulfide absorbs some of the produced sulfur, changing the sulfidation state. Therefore, although iron sulfide catalyzes the reaction and has a high selectivity, it does not act as a true catalyst. For a large-scale process, the changes in the sulfidation state would lead to control problems which are not normally associated with more stable catalysts. For this reason, experiments with FeS were performed a t only two inlet ratios, and attention was focused on the more stable compounds, MoS2 and WS2,which did not show any change in bulk composition. Figures 1 and 2 show that for both MoS2 and WS2, a t a fixed temperature, the selectivity increases with molar feed ratio. It is important to note that replicate experiments have shown that the variation of selectivity is within 54'0 of the mean value shown in the figures. Selectivity for MoS, is much more sensitive to temperature and molar feed ratio than WS2,particularly at feed ratios greater than 1. However, at 750 "C, the sensitivity to the change in the feed ratio is diminished. Decreasing the reaction temperature leads to increased selectivity for MoS2. In fact, at a feed ratio of 2, the selectivities a t 650 and 675 "C are approximately 3-4 times higher than those obtained at 700 "C. In contrast, the relationship between temperature and selectivity of WS2 (Figure 3) shows that there is a local selectivity maximum at 700 "C. This will be discussed later in terms of the rates of the production of sulfur, H2S, and


In addition to high selectivity, an effective catalyst for the reduction of SO2 must also exhibit high catalytic activity as reflected by the rate of SO2consumption. Figures 4 and 5 show the rates of reaction of SO2 for MoS2 and WS2 catalysts, respectively. The highest rates, for both



Figure 4. Effect of temperature on the rate of SO, consumption over MoS, a t different molar feed ratios.



Figure 5. Effect of temperature on the rate of SO, consumption over WS, at different molar feed ratios.

catalysts, are found when the molar feed ratio is 0.5. The effect of molar feed ratio is stronger for WS,. In addition, the rates are generally less sensitive to the molar feed ratio at the lower temperatures. A comparison of the two figures shows that the rates obtained with WS2 are a t least 1.5 times the rates found with MoS,. This is important because the selectivities for the two catalysts are approximately equal a t low molar feed ratios. Although the determination of the rate law of the reactions was not one of the objectives of this investigation, some information concerning the effect of the reactant concentrations on the rates can be obtained from the experimental results. With a differential reactor, the reactions are considered to take place a t the arithmetic mean of the inlet and the outlet concentrations (Massaldi and Maymo, 1968). Because of the small conversion (less than

Ind. Eng. Chem. Res., Vol. 28, No. 7, 1989 929


Feed Ratio

Figure 6. Effect of temperature on the rate of production elemental sulfur over W S pat different molar feed ratios. 3i 3R


n -








Figure 7. Effect of temperature on the rate of production of H,S over WSz at different molar feed ratios.

20'%0),the mean concentration is very close to the inlet concentration which is related to the feed ratio. Figures 4 and 5 show that, for both catalysts, there is little difference between the rates of SO2consumption for the feed ratios of 1.0 and 1.5. Referring to Table I, the only difference in the reaction conditions for these two ratios was the increase in the initial SOz mole fraction from 25% to 36%, while there was no change in the methane mole fraction of 25 '70, This indicates that a change in the SO2 concentration has little effect on the rate. However, for a ratio of 0.5, the mole fraction of methane was increased significantly to 4070, while the SO2 mole fraction was decreased marginally to 20%. In this case, the reaction rate increased by a factor of 2. This implies that the rate of SO2 consumption is dependent on the mole fraction of methane and not that of SOz. For both catalysts, the rates of production of the sulfur-containing products, elemental sulfur, H2S, and COS, essentially follow the same trends that were described above for the consumption of SO2with respect to the molar feed ratio and temperature. Figures 6-8 show the rates of production of these species for WS2 In all cases, at least 90% of the sulfur available from the reacted SO2 appears as elemental sulfur in the product stream. In general, the rate of production of COS is twice the rate of H2S production. The local selectivity maximum at 700 "C (Figure 3) for WS2 can be explained in terms of the rates of sulfur, H2S, and COS. For a ratio of 2.0, for example, between 650 and 675 "C, the rates of production of elemental sulfur and H2S remain essentially constant a t 3.0 X and 6.0 X mol/(m2.s), respectively. The rate of production of COS increases from 1.1 X lo4 to 1.6 X mol/(m2.s), leading

F75 7W Temperature

725 750 "C Figure 8. Effect of temperature on the rate of production of COS over WSz at different molar feed ratios. 550

Table IV. Effect of Temperature and Molar Feed Ratio on C O , Yield (70) SOZ/CH, = 0.5 SOp/CH4 = 1.5 T , "C MoS, WS, FeS MoSl WS, FeS 650 91 61 95 97 72 99 700 55 38 72 88 58 70 750 33 18 35 45 42 60

to a decreasing selectivity. Between 675 and 700 "C, however, selectivity increases. This is due to the fact that, while the rates of production of elemental sulfur and H2S double, the rate of production of COS increases by a factor of only 1.5. Above 700 "C, the selectivity decreases because the relative increase in the rate of production of elemental sulfur with temperature is much smaller than the relative increase in the rates of COS and H2S. Carbon Species Analysis. The catalyst that is used for the reduction of SOz not only affects the production of sulfur-containing products, it also determines the rates of production of carbon-containing species. Carbon dioxide yield ( Y ) is defined as

According to this definition, the carbon dioxide yield represents the percentage of CH4 that has reacted to form the desired carbon-containing product, C02. As with the sulfur selectivity, carbon dioxide yield is essentially a measure of the degree to which the primary reaction takes place. Carbon dioxide yield results are presented in Table IV. In general, for all catalysts, carbon dioxide yield decreases with increasing temperature and decreasing molar feed ratio. For MoS,, the yield at 650 "C is over 90'70. As expected, the lowest values are found at a temperature of 750 "C. These range from 33% to 45%. For WS2, the highest yield is 72%, and it is found at a temperature of 650 "C and a feed of 1.5. At the other extreme, the yield drops to 18% at 750 "C and a ratio of 0.5. The lower values for WS2 in comparison with those of MoS, are due to the higher production of elemental carbon and CO. In the case of CO, production rates were particularly high at the feed ratio of 0.5. Table IV also shows that carbon dioxide yields for iron sulfide are relatively high. In fact, for a feed ratio of 0.5, the yields are similar to MoSz with a feed ratio of 1.5. The final aspect to be considered in the carbon species analysis is the elemental carbon production. The production of elemental carbon was predicted in the thermodynamics analysis of Sarlis and Berk (1988). Carbon

930 Ind. Eng. Chem. Res., Vol. 28, No. 7 , 1989



Figure 9. Effect of temperature on the rate of production of C over MoSz at different molar feed ratios.

Figure 10. Effect of temperature on the rate of production of C over WS2 at different molar feed ratios.

production, via cracking of methane, is a series undesired side reaction since it can lead to rapid deactivation of the catalyst through carbon deposition on its surface. The rates of production of elemental carbon at different temperatures and molar feed ratios for MoS, and WS, are shown in Figures 9 and 10, respectively. As with all other products, for a given molar feed ratio, the rate of carbon production is higher with WSz than with MoS,; however the difference in rates strongly depends on the reaction temperature. At 650 and 675 "C, a minimal quantity of carbon is produced with MoSz. Compared with MoS,, WS2 shows a very different behavior. Even at low temperatures, measurable quantities of carbon were produced. Clearly, MoS, and WS, catalyze the reduction of SOz with CH4 by different mechanisms. Catalyst Comparison. One of the final comparisons to be made is between alumina and the metal sulfide catalysts, MoSz and WS,. Using the dual-site rate expression of Helmstrom and Atwood (1978), which is applicable when no byproducts are formed and the catalyst is activated bauxite, the rate of CH, comsumption was calculated for a composition corresponding to the feed ratio of 2.0 and temperatures of 650 and 675 "C. Under these conditions, no byproducts were formed when MoSz was the catalyst. The calculated rate which was based on the mass of catalyst was converted to a surface area basis by multiplying by 125 m2/g, which is the surface area measured from samples of bauxite catalyst purchased from the same supplier. A comparison showed that the calculated rate for bauxite is twice the experimental rate obtained with MoS,. In order to compare the more general cases when byproducts are formed, the results of Sarlis and Berk (1988)

were used. The comparison was made at 700 "C and a feed ratio of 1.0. At the specified conditions, the SOz consumption rate with alumina and MoSz is similar a t 5.0 X and 5.3 X mol/(m2-s),respectively. However, the rate with WS2 is twice the rates with the other two catalysts a t 10.1 x lo-* mol/(m2.s). The advantages of the metal sulfides become more obvious when the sulfur selectivities are considered. The selectivity for alumina is only 4, whereas the selectivities for MoSz and WSz are significantly higher a t 14 and 13, respectively. These results indicate that both metal sulfides are better catalysts than alumina for the reduction of SOz with CH,. While the activity of MoS, is approximately the same as alumina, the significant increase in sulfur selectivity implies that the rate of production of elemental sulfur is actually 20% higher for MoSz. Comparing WSz and MoS,, it has been shown that the catalytic activity of WSz is almost twice that of MoSz,while their selectivities are almost equivalent. However, the advantage of the latter is more apparent when the carbon dioxide yield and carbon production rates are considered. The carbon dioxide yields for MoS, and WS2 are 69% and 57 % , respectively. More importantly, the carbon production per mole of SO, reduced for MoS, and WSz is 0.18 and 0.25, respectively. This implies that approximately 40% more carbon would be produced per mole of SOz reduced using WS, than with MoSz as the catalyst. Such production of carbon could eventually poison the catalyst, leading to decreased activity. For this reason and the fact that MoS, is less expensive than WSz, it is concluded that MoS, can be considered to be a more effective catalyst than WS,.

Conclusions All metal sulfides tested proved to be catalytic for the reduction of SOz with CH4. Iron sulfide provided high selectivities and carbon dioxide yields but tended to form nonstoichiometric metal sulfides. Both MoS, and WSz were found to be stable under all reaction conditions. MoS, showed high selectivities and SOzconsumption rates equal to those of alumina. In addition, WS2 provided improved selectivities over alumina and SO, consumption rates twice those for alumina, indicating that both MoS, and WSz are better catalysts than alumina in these respects. However, when carbon production results and carbon dioxide yields are considered, it is concluded that MoS, is the better catalyst because of the higher carbon and CO production rates with WS,. Acknowledgment The authors thank the Natural Sciences and Engineering Research Council of Canada for the financial support for this project. Registry No. SOz,7446-09-5; S, 7704-34-9; CH,, 74-82-8; FeS, 1317-37-9; M o S ~1317-33-5; , WS2, 12138-09-9.

Literature Cited Chivers, T.; Hyne, J . B.; and Con, C. The Thermal Decomposition of Hydrogen Sulphide over Transition Metal Sulphides. Int. J . Hydrogen Energy 1980 5 , 499-506. Helmstrom, J. J.; Atwood, G. A. The Kinetics of the Reaction of Sulfur Dioxide with Methane over a Bauxite Catalyst. Ind. Eng. Chem. Process Des. Deu. 1978 17, 114-117. Hunter, W. D.; Wright, J. P. Sulfur Dioxide Converted to Sulfur in Stackgas Cleanup Route. Chem. Eng. 1972, 79, 50-51. Joffre, J.; Geneste, P.; Lerner, D. A. A Quantum-Chemical Study of Site Modelling for Adsorption and Desulfurization of Thiophene. .I. Catal. 1986, 97, 543-548.

I n d . E n g . C h e m . R e s . 1989,28,931-940 Massaldi, H. A.; Maymo, J. A. Error In Handling Finite Conversion Reactor Data by the Differential Method. J. Catal. 1968, 14, 61-68. McMillan, D. Process for Reduction of SOz with Hydrocarbon Vapor. US Patent 3615221, 1971. Mulligan, D. J. Reduction of Sulphur Dioxide Over Transition Metal Sulphides. M. Eng. Thesis, McGill University, Montreal, Canada, 1988.


Rosenberg, H. L.; Engdahl, R. B.; Oxley, J. H.; Genco, J. H. The Status of SOz Control Systems. Chem. Eng. h o g . 1975, 71, 66. Sarlis, J.; Berk, D. Reduction of Sulfur Dioxide with Methane over Activated Alumina. Ind. Eng. Chem. Res. 1988, 27, 1951-54.

Received f o r review August 5, 1988 Revised manuscript received January 23, 1989 Accepted February 19, 1989

High-Temperature Sulfidation-Regeneration of Cu0-A1203 Sorbents V. Patrickt and G. R. Gavalas* Department o f Chemical Engineering, California Institute of Technology, Pasadena, California 91 125

M. Flytzani-Stephanopoulos and K. Jothimurugesan Department o f Chemical Engineering, Massachusetts Institute o f Technology, Cambridge, Massachusetts 02139

Thermogravimetric analysis (TGA) and flow-reactor experiments were used to study sulfidationregeneration of highly porous Cu0-A1203 sorbents. In TGA studies, sulfidation of reduced sorbents produced a high-temperature form of digenite ( C U ~ + ~asS the ~ ) major crystalline product. When a platinum pan was used in the TGA, significant sulfur chemisorption on alumina occurred. Sulfur chemisorption on alumina was eliminated by use of a quartz sample pan. Reaction of the mixed-oxide sorbents with a mixture of H2S, H2, H20, and N2 in a packed-bed microreactor, a t temperatures between 550 and 800 "C, yielded prebreakthrough outlet-H2S levels considerably lower than those predicted by the sulfidation equilibrium of metallic copper. Independent reduction experiments confirmed t h a t alumina stabilizes CuO against complete reduction t o Cu. With this mechanism a t work, low prebreakthrough H2Slevels are attributed to sulfidation reactions of copper a t oxidation states of +1 or +2. T h e sulfided sorbents were completely regenerable in air/N2 mixtures with no deterioration of subsequent sulfidation performance. Among the most promising new technologies of power generation from coal are gasification integrated with combined-cycle power generation (gas turbine in series with a steam turbine), and coal gasification followed by oxidation in a molten carbonate fuel cell. Both technologies require removal of H2S (and COS) from the fuel gas prior to combustion. While this removal can be carried out at ambient temperatures by established technology, removal at high temperatures offers considerable improvement in process economics (Marqueen et al., 1986). In the original research on hot-gas desulfurization, pure metal oxides were tested, including those of zinc and iron (Morgantown Energy Technology Center, 1978). Recent work has focused on mixed metal oxides, especially zinc ferrite (ZnFe,O,). Above about 600 "C, however, sorbents containing zinc oxide lose zinc by reduction to the volatile metal. Iron and copper oxides are not subject to this limitation but do not possess sufficiently large sulfidation equilibrium constants to provide the required level of sulfur removal. In the case of copper oxide, the equilibria 2CuO + H2S + Hz = CuzS + 2Hz0 (1) CU~O + H,S CUZS + HzO (2) are very favorable. In the fuel gas atmosphere, however, the reaction CUO + H2 = CU + H2O (3) progresses rapidly so that sulfidation proceeds by reaction with metallic copper. 2Cu + H2S = CuzS + Hz (4) Present address: CR&DS, Monsanto Company, Creve Coeur,

MO 63167. 0888-5885/89/2628-0931$01.50/0

Table I. Equilibrium Constants for Sulfidation Reactions 1,2. and 4 from Data b s Robie et al. (1984) tema. K 7 00 800 900 1000 1100

reaction 1 21.3 19.2 17.6 16.3 15.2

reaction 2 11.0 9.9 9.0 8.3 7.8

reaction 4 4.1 3.7 3.4 3.2 3.0

As shown in Table I, the equilibrium constant of reaction 4 is much lower than those of reactions 1 and 2. The constant of reaction 4 is such that sulfidation of metallic copper generally does not provide adequate H2S removal for the aforementioned power generation applications. We have recently found that when copper oxide is employed in association with aluminum oxide, or iron oxide, or both, the level of sulfur removal is much higher than that obtained with pure copper oxide (Tamhankar et al., 1986). Such behavior was attributed to the retardation of copper reduction to the metallic form engendered by the association with iron or aluminum oxides. Previous work by Flytzani-Stephanopoulos et al. (1985) on the copper-containing mixed-oxide sorbents involved measurements of breakthrough curves in a packed-bed reactor and X-ray diffraction (XRD) analysis of fresh, sulfided, and regenerated sorbents to determine the crystalline phases present. The purpose of the present study was to obtain detailed mechanistic and qualitative kinetic information about the sulfidation and regeneration of the binary Cu0-A1203 sorbent. To this end, we have used extensively thermogravimetric analysis (in isothermal or temperature-programmed mode) combined with XRD and packed-bed reactor experiments. Some thermo0 1989 American Chemical Society