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
432
Ind. Eng. Chem. Prod. Res. Dev., Vol. 22, No. 3, 1983
Scheme I
100
K,CO, + H,O + CO -+ KHCO, + H C 0 2 K HC0,K + H,O + KHCO, + H, :.K,CO, + 2H,O + CO -+ ZKHCO, + H,
I>
KOH
:.H,O
at > 250 "C at > 250 "C
K,C03
+
90
80
H,O +z
+ CO,
+ CO + H, + CO,
U
70
z 0 U
Scheme I1
60
K,CO, + H,O + 2KOH + CO, 2CO + 2KOH + 2K0,CH 2K0,CH + 2H,O + 2HC02H + 2KOH 2HC0,H -+ 2C0, + 2H, 2KOH + CO, + K 2 C 0 , + H,O :.2H,O + 2CO --f 2C0, + 2H,
catalysis. This mechanism is represented in Scheme 11. The two mechanisms differ in that Yoneda's involves decomposition of formate to produce hydrogen and decomposition of bicarbonate to regenerate carbonate, while Royen and Ehrhard specify formic acid decomposition as the source of hydrogen while carbonate is regenerated by the reaction of hydroxide and carbon dioxide. Zielke et al. (1976) did not distinguish the differences between these two mechanisms since the two essential steps, formate formation and decomposition, are common to both. We report evidence here that suggests that neither accurately describes the chemistry involved in this complex, dynamic system. A key to elucidation of the chemistry is the fate of the formate in the system. As reported by Hartman and Hisatsune (1966),the primary pyrolysis product of sodium formate is sodium carbonate with CO and H2 byproducts and the decomposition takes place at temperatures less than 330 OC (Takagi, 1939). They reference other researchers (Freidlin, 1938; Toyoda, 1950) who suggest that the byproduct gases are derived from a formaldehyde intermediate. Based on our results, we believe this to be the pathway by which sodium formate decomposes in the aqueous system. Experimental Section The pressurized reactor used in our experimentation has been described previously (Elliott and Sealock, 1983). A typical experiment proceeded as follows: water was charged to an autoclave liner, with or without catalysts, and the liner was placed into the autoclave; the autoclave was sealed, purged with nitrogen and then pressurized with carbon monoxide to 500 psig. The autoclave was heated to the reaction temperature where the temperature was stabilized for the length of the experiment. Samples were withdrawn at predetermined intervals through either a gas or liquid sample port by means of a remotely controlled sampling system. At the end of the experiment the autoclave was cooled, vented, and the aqueous product was recovered from the vessel. The use of both a liquid sample port and a gas sample port was necessary in order to monitor the formate/carbonate equilibrium in our system. All gas compositions were determined with a Carle AGC-S gas chromatograph set up for refiery gas analysis. Integration was performed by a Spectra-Physics SP4000 central processor. In addition to gas product analysis, the liquid samples were analyzed by nuclear magnetic resonance spectrometry employing a Varian FT-80 to determine formate concentration by 'H resonance compared to
z U Y 0
50 x
4 I
H
QO
e U U
30 U
9 20
10
0.6w1
A T 300*C 0 10
0
20
30
1 ) uo
50
60
TIME A T TEMPERATURE, MINUTES
Figure 1. Sodium carbonate concentration effect on catalytic conversion of carbon monoxide by the water-gas shift reaction.
0
Na, OXALATE Na2COl
-
0 A
Na OFORMATE H
-
-
0.0 90
80
70
60
50
40
30
20
10
0
REMAINING CARBON MONOXIDE CONCENTRATION, PERCENT
Figure 2. Comparison of activity of several sodium salts as a function of concentration.
an acetate internal standard. Also, liquid samples were dried overnight in an oven at 110 "C and the remaining solid was then analyzed by infrared spectroscopy employing a Perkin-Elmer 283 in a fluorolube mull between potassium bromide disks. The catalysts were reagent grade chemicals used as received from the supplier. Carbon monoxide was CP grade from Matheson (99.5% minimum purity). Water was processed by distillation prior to use. Results Catalyst Concentration Effects. The determination of relative concentration effects of several species involved in a reaction system can provide a clue to the mechanism involved. A study was made of the conversion of carbon monoxide by the water-gas shift reaction over a range of concentrations of aqueous sodium carbonate. The results
Ind. Eng. Chem. Prod. Res. Dev., Vol. 22, No. 3, 1983 433 Table I. Reaction Products from Proposed Mechanistic Intermediaries intermediate conditions reactants sodium bicarbonate sodium bicarbonate sodium formate sodium formate sodium formate sodium hydroxide sodium carbonate sodium carbonate formic acid
300 "C/1500 psig 300 "C/2000 psig 300 "C/lSOO psig 300 * 1 0 "C/l atm 300 "C/lSOO psig 300 "C/1700 psig 300 "C/1500 psig 350 "C/4300 psig 300 "C/1900 psig
CO/H,O
eo
N,/H,O N,/water vapor CO/H,O CO/H,O CO/H,O N,/H,O CO/H,O
are shown in Figure 1. Reproducibility of the data is f2% of carbon monoxide concentration for the final data point. Figure 2 presents the final data point from each curve in Figure 1 as a single curve to further demonstrate this concentration effect. Similar data for sodium hydroxide are also included. The hydroxide curve is distinctly different from the carbonate curve. We interpret these results to show that a twofold concentration increase is required for sodium hydroxide to cause an amount of water-gas shift equivalent to that with sodium carbonate. These results suggest the conversion of one mole of sodium carbonate into two moles of sodium hydroxide as an integral part of the chemical mechanism. The formate data point indicates a slightly better activity than would be expected based on its normal equivalence with hydroxide. The oxalate data also suggest a slightly better than expected activity based on its normal equivalence with sodium carbonate. These points may be explained by the higher solubility of carbon monoxide in formate solutions (Kirk-Othmer, 1963) which may have given the formate and oxalate (which decomposes to formate) an initial advantage in our reactor system. The question of whether the insoluble carbonate catalysts (Elliott and Sealock, 1982) operate by a similar solution catalysis or by a heterogeneous surface catalysis can also be addressed by considering concentration effects. The effect of increasing the amount of the relatively insoluble cadmium carbonate catalyst in the reactor should have no effect on the rate of conversion by solution catalysis since even a minimal amount produces a saturated solution. A concentration effect could only signal a greater amount of reaction due to the larger surface area of the greater amount of slurried powder. The results, as shown in Figure 3, suggest only a negligible difference in conversion over a wide range of concentrations of cadmium carbonate in the aqueous system. However, consideration was not given to mass transfer effects. The appropriate experiments with various stirring speeds were not performed. Therefore, the minor amount of concentration dependence, although only barely greater than experimental error, may be indicative of an alternate, heterogeneous mechanism for the insoluble salts. It appears that although the concentration effect is not strong (based on these experiments), the fact that cadmium carbonate also does not generate sufficient hydroxide ion concentration compared to alkali carbonates strongly suggests that the insoluble transition metal catalysts must be operating by a mechanism somehow different from the basic catalysis represented by alkali and ammonium salts and hydroxides. Pathways for Reaction Intermediates. The behavior of several of the proposed alkali metal intermediates under various conditions and in the presence of various reactants was determined. The results of these experiments, as shown in Table I, suggest the stability of sodium formate in an alkaline solution with a carbon monoxide overpressure even a t 300 "C. We found that bicarbonate, carbonate, and hydroxide were all converted to formate
products
conversion
sodium formate/CO,/H, sodium carbonatelCO,iH, sodium carbonate sodium carbonate/H,/CO sodium formate/H,/CO, sodium formate/H,CO, sodium formate/H,/CO, sodium carbonate/H, (trace) CO,/H,
100%/h 100%/h -40%/h 1%/1.5 h 0%
100%/h 90%/h 0% 100%/h
-
100
90
80 I-
8 70
z
2 U
E z 8
60
0
:
50
X
P
2 E
40
z
5 U zE
30
U
B 20
10
A T 300'C
0 0
10
20
30
40
50
60
TIME AT TEMPERATURE, MINUTES
Figure 3. Concentration effect for cadmium carbonate in the aqueous catalyst system for the water-gas shift reaction.
as determined by IR analysis of the solid recovered by evaporation at 110 "C. This statement also holds for potassium compounds. Formation of formate contrasted with the decomposition of formate to carbonate at these temperatures in the absence of CO (see also Table I), but the rate of this reaction was relatively slow as shown by the low percentages of conversion. As a corollary, sodium carbonate was recovered intact at temperatures up to 350 "C when in the absence of carbon monoxide. Sodium bicarbonate decomposed to sodium carbonate (perhaps via hydroxide) in the absence of water at these temperatures and under no conditions was sodium bicarbonate recovered as a decomposition product from sodium formate. Formic acid, in the presence of CO and water at 300 "C, decomposed to carbon dioxide and hydrogen gas. The Formate/Carbonate Equilibrium. The above results suggest a formate/carbonate equilibrium in this reaction system which may play a crucial role in the mechanism. The thermal stability of various formate species has been studied by a number of research groups (Zapletal et al., 1956; Grant and Turner, 1950;Hisatsune
434
Ind. Eng. Chem. Prod. Res. Dev., Vol. 22, No. 3, 1983
Table 11. Sodium Formate Experimental Results With a CO Atmosphere temp, cat. recovered co "C recoverya solids convC CO,/H,d ._ __ 150 1% l/O 200' 9.7110.9 formate 8% 612 250e 9.0/10.9 formate 13% 81 5 72% 44/28 8.8110.9 formate 300e 76% 39/37 300 9.2110.9 formate 100% 43/57 350e 7.8110.9 carbonatel formate (50:50) 9 1 4 5 % 39/52 to 10.1/10.9 carbonate/ 350f 47/38 formate (75:25) 390e*g 7.4110.9 carbonate/ 100% 38161 formate (40: 60) a Grams of solids recovered by evaporationlg of catalyst initially. Infrared spectrophotometry analysis of 100 - (percent CO remaining in reacrecovered solids. tor). Percent CO,/percent H, in gas phase remaining in reactor. e Experiment sampled through liquid phase port. Results after four batch cycles; ranges indicate changes from first to fourth cycle. 45 min at temperature; all others are 60-min experiments.
and Hartman, 1964). These results show a range of possible reactions from copper formate decomposingto copper metal at around 200 "C, to cadmium and zinc formate decomposing to oxides around 350 "C (although Zapletal et al. (1956) identified cadmium carbonate as a product) to heavy alkaline earth formates decomposing to carbonates at 400-500 "C. Our results show that the transition metal formates (except copper) and alkaline earth formates were converted to carbonates under reaction conditions of 300 "C for 1h in the presence of high-pressure water and carbon monoxide. The copper formate apparently was converted to copper metal as suggested by the literature (Zapletal et al., 1956). These reactions appear to be driven by precipitation of the products from solution. Due to the thermal instability of the sodium formate (Toyoda (1950) reports that the pure compound decomposes at 310 "C), the equilibrium between formate and carbonate is quite temperature sensitive in the range of our study. The catalyst solution approaches an equilibrium concentration at the given temperature as shown by two sets of similar tests, one beginning with sodium formate and one beginning with sodium carbonate. Table I1 shows that the sodium formate remained unchanged at temperatures up to 300 "C. After a single 1-h tests at 350 "C, approximately 50% of the formate had decomposed to sodium carbonate. Following a four-cycle experiment, wherein the catalyst solution was maintained in the reactor and additional carbon monoxide was injected into the reactor for four separate 1-h experiments at 350 "C, the formate had decomposed to carbonate to an extent of approximately 75%. Finally, we see that even in a short experiment at 390 "C approximately 40% of the formate catalyst was converted to carbonate. The differences between gas concentration and conversion for the liquidsampled experiments and the gas-sampled experiments can be traced to the solubility of COz in the aqueous phase. Results of similar experiments with sodium carbonate as the catalyst are shown in Table 111. In this case, we see the beginning of the conversion of carbonate to formate at 200 "C. After 1 h at 250 "C essentially complete conversion has taken place. In the test at 300 "C the conversion was not complete after 1h, but following a fourcycle experiment as described above, only formate was recovered from the catalyst solution. At higher tempera-
Table 111. Sodium Carbonate Experimental Results With a CO Atmosphere* temp, cat. "C recovery' 200 250 250 300 300
recovered solidsb
co
convC
17.5117.0 carbonate 0% (minor formate) 14.5117.0 formate (trace carbonate) 21.8/17.0 formate/carbonate 74% (70:30) 22.811 7.0 formate 83-79%
380
12.7/17.0 carbonate (trace formate) 400f 16.0/17.0 carbonate (trace formate) 400e 11.7117.0 carbonate (trace formate) a See footnotes in Table 11.
91%
CO,/H,' 010
48/27 5113341/38 46/45
100-98% 4315747/51 100% 44/56
tures the equilibrium appears to shift back and only traces of formate were found in the carbonate catalyst solution. These two sets of data consistently support the hypothesis of a formate/carbonate equilibrium in the presence of carbon monoxide wherein formate is the favored species below 300 "C and carbonate is favored above 380 "C. Between 300 and 380 "C lies a region of transition. The Role of Formaldehyde. When formate decomposes to carbonate, carbon monoxide, and hydrogen are typically reported as the coproducts. It has been postulated by Freidlin (1938) that formaldehyde is the precursor of these gases, and formaldehyde has been reported by Toyoda (1950) as a coproduct. Based on a kinetics study of formate decomposition Hartman and Hisatsune (1966) proposed a mechanism which specifies a transition complex formed from two formate ions that decomposes to carbonate and formaldehyde.
-
0
I
H
In our study we have not attempted to monitor formaldehyde production; however, we have produced evidence which suggests the presence of formaldehyde in the aqueous product solution in some cases. In order to follow the formation and decomposition of formate in our reactor system, liquid samples were withdrawn and analyzed by proton NMR. These analyses gave an inconsistent picture of formate concentration, but the formaldehyde peaks were obliterated by the water peak. There were clearly some cases in which no formate was detected while evaporation of the same sample yielded a white powder which was determined to be formate by IR analysis. This finding suggested that if all the formate had decomposed in the experiment as shown by the liquid sample, then dissolved formaldehyde might be reacting with the carbonate to reform formate during the heating cycle of evaporation. We have now shown that this reaction could be the explanation for the presence of formate. Sodium carbonate can be reacted with formalin solution at temperatures from 80 to 120 "C to produce sodium formate, methanol, and carbon dioxide by the Cannizzaro reaction (Elliott et al., 1982). These experiments demonstrated that formalin solutions of sodium carbonate, when heated to dryness, yielded only sodium formate and when dried at room temperature resulted in no reaction, so that formaldehyde and sodium carbonate were recovered as
Ind. Eng. Chem. Prod. Res. Dev., Vol. 22, No. 3, 1983 435
shown by infrared analysis. The implication of formaldehyde in our reaction system supports the mechanistic contentions of Hartman and Hisatsune (1966). Conclusions The elucidation of formate as an important intermediate in alkali catalysis of the water-gas shift reaction is generally credited to Yoneda et al. (1943a,b; 1944a,b). His study of aqueous potassium carbonate solution's effect on carbon monoxide led him to postulate the mechanism described earlier in this paper. We now believe that this series of reactions inadequately portrays the chemistry involved in this system in that the importance of bicarbonate is overemphasized. We see no evidence for the formation of bicarbonate as the byproduct from the reaction of formate with water as proposed by Yoneda as the hydrogen production step. While our experiments verify that bicarbonate decomposed to carbonate in the absence of water, the carbonate is quickly converted to formate in the presence of both water and CO. In a similar vein, although we have verified that formic acid does decompose to hydrogen and carbon dioxide under the conditions of our system, direct decomposition of sodium formate is simpler and equally credible compared to the pathway suggested by Royen and Ehrhard (1953 and 1956). Therefore, we propose the following as a depiction of the chemistry involved in the water-gas shift reaction in the aqueous catalytic system.
In this system, sodium carbonate is driven to complete reaction with water by the simultaneous and fast reaction of carbon monoxide with the carbonate-generated hydroxide ion to produce the formate. The formate thermally decomposes to sodium carbonate and formaldehyde as suggested by Hartman and Hisatsune (1966). The formaldehyde is the source of hydrogen as produced in the water-gas shift. In this mechanism, bicarbonate exists only as the intermediate ionization step in the reaction of carbonate and water. With the strong driving force of CO reacting with hydroxide, the bicarbonate is quickly reacted
to hydroxide to push the system toward equilibrium. This mechanism suggests that any hydroxide ion generator can act as a water-gas shift catalyst and explains the activity of all the acid salts tested thus far. The importance of the stability of the particular formate salt in the system is also suggested since its decomposition appears to be the rate-limiting step. The low temperature of decomposition for the alkali formates to carbonates compared to the higher temperature required for the alkaline earths may be an explanation for their differences in catalytic activity (Elliott and Sealock, 1982). This difference may also explain why some of the transition metal salts are more active than the alkaline earths even though they are both nearly insoluble and generate a relatively low concentration of hydroxide. 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. Carbonate, 3812-32-6; hydroxide, 14280-30-9; formate, 71-47-6.
Literature Cited Elliott, D. C.; Sealock, L. J. Jr. Ind. fng. Chem. prod. Res. Dev. 1983. preceding article in this issue. Elliott, D. C.; Sealock, L. J. Jr.; Hallen, R. T., submitted for publication in Chem. Commun., 1983. Freidiln, L. Kh. J . Appl. Chem. USSR 1938, 11, 975. Grant, D. 0.; Turner, H. S. Natwe (London), 1950, 165, No. 4187. 153-4. Hartman, K. 0.; Hlsatsune, I. C. J . phvs. Chem. 1988, 7 0 , 1281-1287. Hlsatsune, I. C.; Hartman, K. 0. Sclence 1984, 145, 1455-6. "Klrk-Othmer Encyclopedia of Chemlcai Technology", 2nd ed.; Wiley: New York, 1963. Royen, P.; Ehrhard. F. Erdol Kohk, 1953, 6 , 195-9. Royen, P.; Ehrhard, F. ErdOKohh9 1958, 9 , 19-24. Takagl, S. J . Chem. Soc. Jpn. 1939, 60, 625; Chem. Abstr. 1942, 3 6 , 6401. Toyoda, R. Bull. Inst. Chem. Res. Kyoto Unlv. 1950, 20, 11. Yoneda, K.; Honda, Y.; Momiyana, N.: Abe, R. J . Chem. Soc.Jpn. 1943a, 46, 554. Yoneda, K.; Kondo, S.; Abe, R. J . Chem. SOC.Jpn. 1943b, 46, 667. Yoneda, K.;Kondo, S.; Abe, R. J . Chem. Soc. Jpn. I941a, 44, 385. Yoneda, K.; Kondo, S.; Abe, R. J . Chem. Soc. Jpn. 1941b. 4 4 , 388. Yoneda. K.; Kondo, S.; A k , R. J . Chem. SOC.Jpn. 1944a, 4 7 , 5 . Yoneda, K.; Kondo, S.; Abe, R. J . Chem. SOC. Jpn. 1944b, 4 7 , 7. Zapletal, V.; Jedllcka, J.; Ruzkka, V. Chem. Lbty 1958, 50, 1406-9 Chem. Abstr. 1959. 51, 2438. Zlelke, C. W.; Rosenhoover, W. A.; Gorin. E. Prepr., Dlv. Fuel Chem., Am. Chem. SOC.1978, 21(7). 163-186.
Received for review October 1, 1982 Accepted February 17, 1983
