426
Ind. Eng. Chem. Prod. Res. Dev. 1983, 22, 426-431
laboratory. We hope to publish the results of tests on these catalysts in the near future.
Hopkins, P. D.; Meyers, E. L. Amoco oil co., unpublished data, 1978. Houalla, M.; Delannay, F.; Delmon, 8 . J. Phys. Chem. 1981, 85,
Registry No, Co, 7440-48-4; Ni, 7440-02-0; Mo, 7439-98-7; c o ~ &13762-146; ,, N ~ M ~ 14177-550; o~, M ~ ( M ~15123-80-5, ~ ~ ) ~ ,
Johnson, M. F. L.; Erlckson, H. US. Patent 4 202 798, 1980. Kana c. c.; Johanson, E. S. "LlquM Fuels from Coal"; Elllngton, R. T. Ed.; Academic Press: New York, 1977;p 89. kine, J.; Pratt, K. C. Ind. Eng. Chem. Fundam. 1981, 20, 1-5. Massoth, F. E. A&. Catal. 1978, 27, 265-310. Nahin, P. G.;Huffman, H. C. Ind. Eng. Chem. 1949, 4 7 , 2021-2026. Parekh, E. S.;Weller, S. W. J. Catel. 1978, 55, 58-62. Ratnasamy, P.; Mehrotra, R. P.; Ramaswamy, A. V. J. Catal. 1974, 32,
L i t e r a t u r e Cited Brown, S. M.; Wallace, D. S. US. Patent 3969273,1976. Choca, M. E. U.S. Patent 4066572, 1978. Colgan. J. D.; Chomitz, N. US. Patent 3 287 280, 1966. Dufresne, P.; Payen, E.; Grimblot, J.; Bonnelle, J. P. J. Phys. Chem. 1981,
85,2344-2351. Gajardo, P.; Grange, P.; Delmon, E. J. CataI. 1980, 63, 201-216. Qates, B. c.; Katzer, J. R.; Schuh, G, c. A, "Chemistry of Catalytic Processes"; McGraw-HIII: New Yo&, 1979;p 414. Glordano, N.; Bart, J. C. J.; Vaghi, A.; Castellan, A.; Martinotti, G. J. Catal. 1975, 36, 81-91. Greegor, R. B.; Lytle, F. W.; Chin, R. L.; Hercules, D. M. J. phvs. Chem. 1981, 85,1232-1235.
1704-1709.
63-71. Rigge, R. J.; Olsen, 0. C.; Goetschel, C. T. US. Patent 3879310, 1975. Wllllams, A.; Butler. G. A.; Mmmonds, J. J. Catal. 1972, 24, 352-355. Young, D. J.; Udal& P.; Trimm, D. L. "Proceedings International Symposium Catalyst Deactivation-Antwerp, 1980";Delmon, E.; Froment, G. F.;Ed.; Elsevier: New York, 1980;pp 331-340.
Received for reuiew September 8, 1982 Accepted February 22, 1983
Aqueous Catalyst Systems for the Water-Gas Shift Reaction. 1. Comparative Catalyst Studies Douglas C. Elliott' and L. John Sealock, Jr. Battelle Pacific Northwest Laboratories, Richland, Washlngton 99352
A comparative study of catalytic material for the water-gas shift reaction has been conducted in a pressurized aqueous system. Resutts of this Investigationhave shown that alkali metal salts catalyze the water-gas shii reaction in the temperature range of 200-400 'C. Alkaline earth salts were found to be significantly less active. Transition metal salts ranged in activity from better than the alkali metals in some cases to nearly as poor as the alkaline earths in other cases. The use of ammonla and ammonium compounds was also demonstrated. Experimental results show that dissolved, alkali catalysts compare favorably with slurried, conventional, supported metal oxide catalysts in the pressurized aqueous system.
Introduction
The water-gas shift reaction involves the reaction of carbon monoxide and steam to produce hydrogen and carbon dioxide and derived its name from its industrial use to increase the hydrogen content of water gas produced from the reaction of steam with hot coke or coal. Current interest in this reaction lies in the tailoring of the hydrogen to carbon monoxide ratio of gas streams employed for chemical synthesis. For example, product gas from a typical coal gasification unit may have an Hzto CO ratio of approximately 1to 1while the synthesis gas composition required for methanol production would have a ratio of 2 to 1 and the ratio for methane synthesis is 3 to 1. The study of catalysts for this reaction has focused primarily on heterogeneous catalysts (Hawker, 1982). Specifically, iron oxide-chromium oxide catalysts have been used at high temperature (350 to 450 OC) (Ruthven, 1969; Podolski and Kim, 1974; Bohlbro, 1966) and newer low temperature (200-260 "C) copperzinc oxide catalysts have been developed for secondary treatment of gas product streams (Dienes, 1967; Reitmeir and Fleming, 1968; Reitmeir, 1968,Gutman and Johnson, 1970; Chandra et al., 1972; Mukherjee et al., 1976; McDonald, 1969). The low-temperature catalyst is particularly susceptible to poisoning by chlorine or sulfur as well as deactivation due to sintering. More recent studies of homogeneous catalysts for the water-gas shift reaction have dealt primarily with 0196-4321/83/1222-0426$01.50/0
organometallics (Cheng and Eisenberg, 1978; Singleton et al., 1979; Ungerman et al., 1979; Baker et al., 1980; King et al., 1980, 1981). The use of a high-pressure water system for the watergas shift reaction was first proposed by Casale (1932) although he was unaware of the potential of basic catalyst solutions. The pressurized water system finds its advantage in a kinetic effect due to pressure as well as a shift in product composition due to the large excess of water driving the reaction to completion. The work of Yoneda et al. (1941a,b, 1943a,b, 1944a,b) during World War I1 established the concept of using an aqueous solution of metal carbonate as a catalyst for the water-gas shift reaction. This work included a comparative study of a large number of transition metals; however, potassium was the only alkali metal tested. The application of this mechanism in a process scheme was recently examined (Zielke et al., 1976) while patents for the use of cesium carbonates and other salts have also been issued (Aldridge, 1971, 1973). Our work reported here goes beyond these previous reports to test and compare additional related compounds for catalytic activity in a pressurized aqueous reaction system. This study has led to a broader definition of an aqueous catalyst for the water-gas shift reaction. In addition, comparisons among the catalysts are made, and activity relationships to commercial supported metal/ 0 1983 American
Chemical Society
Ind. Eng. Chem. Prod. Res. Dev., Vol. 22,
AIR SIGNAL FROM VALVE CONTROLLER
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Figure 2. Sampling bulb system.
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Figure 1. One-liter autoclave modified for multiple remote sampling.
metal oxide catalysts are examined.
Experimental Section A pressurized aqueous batch reactor system was developed wherein gas reactions were catalyzed at high pressure (up to 5O00 psig). The system had the capability for sampling several times during an experiment while the reactor was maintained at reaction temperature and pressure. Since we were studying a saturated aqueous system where there was a liquid water phase and a discrete gaseous phase (when operated below the critical temperature), the reactor was designed to allow samples to be drawn from near the bottom and near the top of the reactor. A drawing of the reactor is presented in Figure 1. The pressure vessel, manufactured by Autoclave Engineers, was a 1-L autoclave equipped with a magnetic-coupled stirrer. The autoclave was operated with standard heating, cooling, and stirring equipment provided with the unit, except the stirring baffle was not used in order to fit a stainless steel liner into the reactor. The use of the liner facilitated product recovery and reactor cleanup, and prevented direct contact of the pressure vessel with aqueous alkalies. The material of construction was Inconel for all reactor parts exposed to the reaction environment. The multiple sampling system used a body wall port for access to the top of the reactor for gas sampling. The standard l/s-in. dip tube sample system was capped and a lI4-in.dip tube inserted through the solids injection port in the cover of the autoclave to allow removal of liquid from the bottom of the reactor. The autoclave and sample removal system were operated remotely after the reactor had been batch charged. As shown in Figure 2, the sample bulb system was isolated from the reactor system by an air-controlled, high-pressure needle valve. When the reactor reached the proper conditions and a sample was required, the air-controlled valve was opened and the sample flowed through the two open valves into the sample bulb, which was monitored via the pressure gauge. After the sample was withdrawn, which required approximately 1s, the valves were closed and the sample bulb was disconnected at the tubing union and another sample bulb connected. The samples were typically withdrawn every
10 or 15 min during experiments which lasted 30 to 60 min at reaction temperature. In this reactor system where a portion of the reaction mixture was withdrawn from the system, the perturbation of that withdrawal and its effect on later samples had to be considered. The effect of the sample removal was measured by replicating the experiments and sampling at different times during the experiments. The variances due to the sample removal were qualitatively indistinguishable from the variances due to experimental error. The sample bulb had an average volume of 111mL and after sampling contained a sample of 10 to 40 mL of condensate at up to 650 psig of gaseous product pressure at ambient temperature. Generally, five samples were drawn during an experiment and pressure decreased with the drawing of each sample. Mass balances in excess of 95% were typical. A representative experiment proceeded as follows: water was charged to the autoclave liner, with or without catalyst, and the liner was placed in the autoclave; the autoclave was sealed, purged with nitrogen, and then pressurized with the reactant gas to a specified pressure. The autoclave was then heated to the reaction temperature where the temperature was stabilized for the length of the experiment. Samples were withdrawn at predetermined intervals through either of the sample ports. At the end of the experiment the autoclave was cooled, vented, and the aqueous product recovered from the vessel. Analysis of the products was performed primarily by gas chromatography. Fixed gases were sampled at ambient conditions by venting the high-pressure bomb through a glass gas sampling bulb with a side arm septum port. A syringe sample of the gas was then removed through the septum and injected into the gas chromatograph. The gas analysis was performed in a Carle AGC-S gas chromatograph set up for refinery gas analysis. The liquid product was poured from the sample bulb after the gas was vented. The pH of the liquids was determined with a pH meter. The solid catalysts were recovered by drying and then analyzed by infrared spectroscopy. The catalysts were reagent grade materials obtained from a number of different suppliers; no attempt was made to repurify them. Carbon monoxide was CP grade from Matheson (99.5% minimum purity). The water was processed by distillation prior to the experiment.
Results Metal Carbonate Catalyst Comparisons. The catalytic properties relative to the water-gas shift reaction of metal carbonates were tested in our multiple remote sampling autoclave system. Experiments were typically performed using a 0.32 M aqueous solution (500 g of distilled water) and with 500 psig carbon monoxide initial pressure. Figure 3 summarizesthe results of these experiments. The graph shows the residual carbon monoxide concentration as mole percent in the reactor as a function of reaction time. The decrease in concentration is a result both of conversion of CO to C 0 2 and also dilution of unreacted CO
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Figure 3. Catalysis of the water-gas shift reaction by various metal carbonate compounds.
with additional hydrogen product gas, the two coincidental effects of the water-gas shift reaction. Therefore, the data accentuates the differences in catalytic activity among the various catalysts and facilitates the comparison of relative activities. However, the data format is not conducive to pure kinetic calculations, i.e., rate constants. The rate of CO disappearance was chosen as the measure of the water-gas shift reaction rate in preference to the rate of C 0 2 or H2 increase. The lack of information on the behavior of C 0 2 over aqueous alkali systems would obscure conclusions based on COz. Our experience has shown that loss of integrity of "grab samples" (as employed in this work) due to leakage on diffusion lessens the reliability H2determinations. However, our results do show that the water-gas shift stoichiometric relations between CO, COz, and H2 were within the range of uncertainty caused by these factors. Even a summary view of Figure 3 indicates a difference between the reactivities with alkali metals vs. reactivities with alkaline earths. It is of interest that some transition metals also exhibit water-gas shift catalytic activity. The nickel carbonate is noteworthy because it also catalyzed methane formation. These results for potassium and cadmium confirm those of Yoneda et al. (1941b, 1943a,b); however, Yoneda's zinc carbonate, in contrast to ours, was nearly as active as his cadmium carbonate. The alkali metals tested exhibit relatively similar activities, and an ordering based on the periodic table of elements is apparently not followed. This is also interesting considering the patent of Aldridge (1971) on cesium catalysts for the water-gas shift reaction. This patent contends that when used on a high surface area support, cesium carbonate exhibits a catalytic activity of up to four times greater than
potassium carbonate on a similar support. The examples given in the patent were performed at temperatures of 330 and 385 "C, which are similar to our own tests. Apparently, there is some dramatic difference in the reactivities of the supported catalysts of the patent vs. the aqueous solution catalysts used in our work. A point of contention which arises when comparing alkali compounds to alkaline earth compounds is that alkalis are typically in a +1oxidation state while alkaline earths are typically in a +2 oxidation state. This means that when comparing equimolar solutions, as was done here, there are only half as many alkaline earth metal ions in the reactor compared to the number of alkali metal ions. Depending on the mechanism of the reaction, this difference might result in lower activities for the alkaline earths. However, when the sodium carbonate concentration is cut to one-half of the original, the reactivity is still significantly greater than any of the alkaline earth carbonates. In addition, transition metal carbonates in +2 oxidation states were also tested; of these carbonates only zinc carbonate did not exhibit nearly as good or better reactivity than the alkali metals. Apparently, the reduced number of metal ions for these metal carbonates does not result in such poor catalytic activity due to the +2 oxidation as for the alkaline earths. A second major difference between the alkalis and alkaline earths is the extent of solubility in water. The limited solubility of the catalyst could, of course, limit the number of active species in solution. The alkali carbonates are typically quite soluble to very soluble, and the alkaline earth carbonates have low, but measurable, solubilities. One would reasonably expect in our reactor system that a water-soluble catalyst might be more reactive than an insoluble catalyst. One could envision the insoluble catalyst lying on the bottom of the reactor separated from the gas phase by the dilute aqueous solution, especially if there was inadequate stirring. However, the transition metal carbonates also have very low solubilities in water yet also show good activity as do the more soluble alkali metal carbonates. Therefore, by analogy, the low solubilities of the alkaline earths should not be the sole factor limiting their activity. A third typical difference between alkali metal and alkaline earth carbonates is in their relative strengths of basicity. At the concentrations used, the alkali metal carbonates generated a pH of 11.5 to 12.0 while the alkaline earth carbonates were at about 9.0. This difference in base strength could affect relative reactivities; however, the transition metal carbonates exhibit even lower pH, about 8.5, yet are nearly equal in activity to the alkali metal carbonates. Three of the major differences between alkali metal carbonates and alkaline earth carbonates have been considered as potential explanations for the difference in reactivity. All three could tend to cause the reduced level of catalytic activity of the alkaline earth carbonates. In contrast, some transition metal carbonates which exhibit all three of these potentially differentiating and detrimental characteristics show reactivities similar to the alkali metals and much greater than the alkaline earths. Further consideration of the catalytic mechanism is given in a companion article (see Elliott et al., 1983). Other Catalytic Aqueous Solutions. Seventeen sodium compounds have been studied for catalytic activity for the water-gas shift reaction with the results presented in Figures 4 and 5. Although several of these anions have been suggested as catalysts, (Zielke et al., 1976) this work represents the first comparative study and the first study
Ind. Eng. Chem. Prod. Res. Dev., Vol. 22, No. 3, 1983 429 100
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of the sodium compounds. A number of similar potassium salts also exhibited activity as shown in Figure 6. The compounds exhibit a range of activities at equimolar
concentration (based on the anion) which may be explainable by either or both of two resulting situations: (1) the equimolar concentrations result in nonequivalent concentrations of metal ion in the reactor; and (2) the various acids have differing ionization constants which affect the amount of hydroxide ion generated in the solutions. Basic solution concentration effects will be considered further in our next article (Elliott et al., 1983). As a sidelight, we have tested the acid forms of two of our catalysts and found that neither formic acid nor carbon dioxide exhibits catalytic activity for conversion of carbon monoxide by the water-gas shift reaction. However, hydrogen and carbon dioxide are produced in the presence of formic acid but only through decomposition of the acid at the reaction conditions. Recovery of the alkali metal salt catalysts in their original form was not accomplished in many cases, e.g., oxalates, phenolate, pyruvate carbonates, and hydroxide. These salts generated formate in this system either through direct decomposition or formation of organic acid and conversion of hydroxide to formate by carbon monoxide. Malonate, propionate, citrates, and tartrates underwent this conversion to a lesser degree. Decomposition of the salts into a catalytic form, i.e., in situ generation of the catalyst, is the likely explanation of the activity in these cases. A catalytic mechanism which incorporates these findings and is based on the decomposition and dissociation of these catalytic materials has been proposed (Elliott et al., 1983). As a result of the differences in catalytic activity which were demonstrated among the various metal cations in these aqueous systems, tests were made to determine if the metal was indeed necessary to the system. Ammonium carbonate and hydroxide, as well as ammonia gas, were tested and the results are seen in Figure 7. The experi-
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ments demonstrate that basic ammonium compounds have strong catalytic activity similar to alkali metal compounds. Therefore, the metal cation, per se, is not a necessary part of the catalytic mechanism.
Supported Metal Oxide Catalysts. Tests have been made with commercial water-gas shift catalysts in order to provide a point for comparison with our aqueous system. In these tests the catalysts were pellets of metal oxides on inert supports as described in Table I. These catalysts were crushed and ground to powder. This powder was then slurried in the water feed at 1 % concentration and placed in the autoclave. The tests were run from that point in a manner similar to those with the solution catalysts and the results are shown in Figure 8. One experiment is shown wherein the pellets were not ground to a powder. Based on comparisons with this one experiment, the limited mass transfer of gas to the solid catalyst is evident. This finding does not imply that the gringing of the catalyst eliminated mass transfer constraints. The temperature effect is evident for Cu-ZnO catalyst; a similar set of curves for the sodium carbonate catalyst is shown in Figure 9. The data, as reported here, suggest that the sodium carbonate solution is at least a comparable
Ind. Eng. Chem. Prod. Res. Dev. 1903, 22, 431-435
catalyst to the metal catalysts examined. In fact, it shows a higher activity at a given temperature in our reactor configuration.
Conclusions We have completed a comparative study of catalytic materials for the water-gas shift reaction. The pressurized aqueous reactor system has been tested with the use of both soluble and insoluble catalysts. Alkali metal salts and some transition metals salts are demonstrably more active catalysts than similar alkaline earth compounds. A wide variety of alkali compounds can act as catalysts (or generate a catalytic species in situ) for the water-gas shift reaction in a pressurized aqueous system. In addition, we have demonstrated that the cation need not be a metal by showing the catalytic activity of ammonium compounds. Comparative tests with supported metal oxide catalysts suggest that the aqueous solutions of sodium carbonate may have similar or increased catalytic activity levels in a pressurized system.
Acknowledgment The authors wish to thank Dave Ortiz, who operated the reactor system for these experiments. This work was funded through the Basic Research Department of the Gas Research Institute (GRI), and the authors thank John L. Cox and Charles R. Nelson of GRI for their support and encouragement of our research. Registry No. CdC03, 513-78-0; NiC03, 3333-67-3; CuC03, 1184-64-1;Ag&O3,534-16-7;C S ~ C O534-17-8; ~, KzC03,584-08-7; Li2C03, 554-13-2; Na2C03, 497-19-8; NH3, 7664-41-7; NH,OH, 1336-21-6;(NH&C03, 506-87-6;ZnC03, 3486-35-9; BaC03, 51377-9; SrC03, 1633-05-2;CaC03, 471-34-1; MgC03, 546-93-0;hydroxide, 14280-30-9; formate, 71-47-6;bicarbonate, 71-52-3;acetate, 71-50-1; carbonate, 3812-32-6;oxalate, 338-70-5; citrate, 126-44-3; hydrogen oxalate, 920-52-5;dihydrogen citrate, 55465-68-4;glyoxalate, 430-75-1; phenolate, 3229-70-7; propionate, 72-03-7; malonate, 156-80-9; ketomalonate, 4004-36-8; K2tartrate, 921-53-9;
43 1
KNa tartrate, 304-59-6; KH tartrate, 868-14-4; KH carbonate, 298-14-6; K quadroxalate, 127-96-8; KH oxalate, 127-95-7.
Literature Cited Aldridge, C. L. U S . Patent 3815218, 1971. Aldrldge, C. L. U.S. Patent 3755558, 1973. Baker, E. C.; Hendricksen, D. C.; Elsenberg, R. J . Am. Chem. SOC. 1980, 102, 1020-1027. Bohlbro, H. “An Investigatlon on the Klnetlcs of the Conversion of Carbon Monoxide wtth Water Vapor over Iron Oxide Based Catalysts”; Gjellerup, Copenhagen, Denmark, 1968. Casale, L. US. Patent 1 843 540, 1932. Chandra. M.; Sign, S. S.; Ghosai, S. R. Ind&n Chem. Eng. 1972, 74(4), 27-33. Cheng, C-H.; Elsenberg, R. J . Am. Chem. Sac. 1978, 700, 5988-5970. Dienes, E. K. U S . Patent 3303001, 1987. Elliott, D. C.; Sealock, Jr., L. J.; Hallen, R. T. Ind. Eng. Chem. Prod. Res. D e v . 1983, following paper in this issue. Qutmann. W. R.; Johnson, R. E. U S . Patent 3548 140, 1970. Hawker, P. N. &jmcarbon Roc8ss. 1982, 67(4). 183-187. King, A. D., Jr.; King, R. B.; Yang, D. C. J . Am. Chem. SOC. 1980, 702, 1028-1032. King, A. D.. Jr.; King, R. B.; Yang, D. C. J . Am. Chem. SOC. 1981, 703, 2899-2704. MoDonaki, W. T. Ph.D. Dissertation, Department of Chemlcai Engineering, University of Loulsvliie, Loulsvllle, KY, 1989. Mukherjee, P. N.; Basu, P. K.; Roy, S. K.; Chatterjee, S. K. Indlen J . Techno/. 1978, 74, 138-141. Podoiskl, W. F.; Kim, Y. 0. Ind. Eng. Chem. Process D e s . D e v . 1974, 73, 415-421. Reltmer, R. E. U.S. Patent 3390 102, 1988. Reitmer, R. E.; Fleming, H. W. U.S. Patent 3338972, 1988. Ruthven, D. M. Can. J . Chem. Eng. 1989, 4 7 , 327-331. Singleton, T. C.; Park, L. J.; Price, J. L. “Water-Gas Shift Reaction in the Homogeneous Rhodium-Iodlne Catalyst System”; presented at the 177th Natlonal Meeting of the American Chemlcal Society, Divlsion of Petroleum Chemlstty, Honolulu, HI, Apr 1-8, 1979. Ungermann, C.; Landis, V.; Moye, S. A.; Cohen, H.; Walker, H.; Pearson, R. 0.;Rinker, R. Q.; Ford, P. C. J . Am. Chem. Soc. 1979, 707, 5922-5929. Yoneda, K.; Kondo, S.; Abe, R. J . Chem. Soc. Jpn. 1941a, 44, 385. Yoneda, K.; Kondo, S.; Abe, R. J . Chem. Soc.Jpn. 194lb, 44, 388. Yoneda, K.; Honda. Y.; Monlyama, N.; Abe. R. J . Chem. Soc. Jpn. 1943a 46. 554. Yoneda, K.; Kondo, S.; Abe, R. d . Chem. SOC.Jpn. 1943b, 46. 887. Zielke, C. W.; Rosenhoover, W. A.; Oorln, E. Prepr. Dlv. Fuel Chem.. Am. Chem. Soc. 1978, 27(7), 183-188.
Receiued for reuiew October 1, 1982 Accepted February 17, 1983
Aqueous Catalyst Systems for the Water-Gas Shift Reaction. 2. Mechanism of Basic Catalysis Douglas C. Elllott,’ Richard T. Hallen, and L. John Sealock, Jr. Battelle Paclflc Northwest Laboratories, Rlchland, Washington 99352
A mechanism is proposed for catalysis of the water-gas shift reaction by basic aqueous solutions in a pressurized
reactor. The mechanism is proposed as a cyclic system with three intermedlate anions: carbonate, hydroxide, and formate. A carbonate ion reacts with water to form hydroxide and carbon dioxide. The hydroxide ion reacts with carbon monoxide to generate formate ion. The formate decomposes to carbonate and formaldehyde which in turn generates hydrogen by decomposition. This mechanism c a n proceed in the presence of any basic catalyst precursor which generates one of the three catalytic species. A study of the formate/carbonate equilibrium was made, as well as catalyst concentration studles.
Introduction Catalysis of the water-gas shift reaction by aqueous solutions of basic compounds has been reported previously (Yoneda et al., 1941a,b; 1943a,b; 1944a,b; Zielke et al., 1976; Elliott and Sealock, 1982). A pressurized system is required for this type of solution catalysis when operated in the temperature range of 2W400 OC. During World War
II (Yoneda et al., 1941a,b, 1943a,b, 1944a,b) first elucidated the importance of the formate intermediary in this process. They developed and investigated a mechanism based on potassium carbonate solution chemistry which is represented by the series of equations in Scheme I. A similar mechanism was also developed by Royen and Ehrhard (1953,1956) based on dry potassium carbonate
0196-4321/83/1222-Q431$Q~.5Q/O 0 1983 American Chemical Society
