Ind. Eng. Chem. Res. 1987, 26, 874-877
874
services. We thank the Senate Research Council, University of the Witwatersrand, Sasol Technology, and the Foundation for Research Development, CSIR, Pretoria for financial assistance. Registry No. Co, 7440-48-4; MnO, 1344-43-0; Fe, 7439-89-6; CO, 630-08-0; CZH4, 74-85-1; CSHs, 115-07-1; CPHfi, 74-84-0.
Literature Cited Barrault, J.; Renard, C. Nouv. J . Chem. 1983a, 7, 149. Barrault, J.; Forquy, C.; Pericheu, V. Appl. Catal. 198313, 5, 199. Barrault, J.; Renard, C.; Yu, L. T.; Gal, J. Proc. 8th Int. Congr. Catal. Berlin 1984, 2, 101. Barrault, J.; Renard, C. Appl. Catal. 1985, 14, 133. Bussemeier, B.; Frohning, C. D.; Horn, G.; Kluy, W. Deutches Offen, 2 518964, 1976. Copperthwaite, R. G.; Hack, H.; Hutchings, G. J.; Sellschop, J. P. F., University of the Witwatersrand, South Africa, unpublished results, 1985a. Copperthwaite, R. G.; Hack, H.; Hutchings, G. J., Sellschop, J. P. F. Surf. Sci. 1985b, 164, L827-L830. Cornils, B.; Frohning, C. D.; Morow, K. Proc. 8th Int. Congr. Catal. Berlin 1984, 2, 23.
Dowden, D. A.; Mckenzie, N.; Trapnell, B. M. W. Proc. R. SOC. London, Ser. A 1956, A237, 245. Friedel, R. A.; Anderson, R. B. J . Am. Chem. SOC. 1950, 72, 1212, 2307. Harrison, D. L.; Nicholls, D.; Steiner, H. J. Catal. 1967, 7, 359. Henrici-Olive, G.; Olive, S. Angew. Chem., Int. Ed. Engl. 1976, 15, 136. Hutchings, G. J.; Boeyens, J. C. A. J. Catal. 1986, 200, 507. Kolbel, H.; Tillmetz, K. D. Deutches Offen. 2 507 647, 1976. Korf, C. J.; Espinoza, R. L. CSIR (CERG) Report 584, 1986; CSIR, Pretoria, South Africa. Kugler, E. L. Prepr.-Am. Chem. SOC., Diu. Pet. Chem. 1980, 564. Maiti, G. C.; Malessa, R.; Baerns, M. Appl. Cata2. 1983, 5, 151. Schulz, H. C1Mol. Chem. 1985, I , 231. Schulz, H.; Gokcebay, H. Catalysis of Organic Reactions; Kosak, J. R., Ed.; Marcel Dekker: New York, 1984; p 153. van der Riet, M.; Hutchings G. J.; Copperthwaite, R. G. J. Chem. SOC.,Chem. Commun. 1986, 788. Varma R. L.; Dan Chu, L.; Mathews, J. F.; Bakshi, N. N. Can. J . Chem. Eng. 1985,63, 72.
Received for review February 24, 1986 Revised manuscript received September 24, 1986 Accepted January 5, 1987
Competition for Active Sites between Hydrogen and Methyl Esters of Fatty Acids in Vapor Phase on a-A1203-SupportedCopper and Nickel Catalysts Jeppe Magnusson Department of Chemical Reaction Engineering, Chalmers University of Technology, S-412 96 Goteborg, Sweden
The H2/D2exchange reaction, catalyzed by the copper and nickel catalysts, was shown to be a model reaction for studying the competition for active sites on the catalyst surface. It was found that hydrogen and the methyl esters of fatty acids compete mutually for the active sites on the supported nickel catalyst. On the copper catalyst, however, hydrogen and the methyl esters compete only for a minor part of the active sites on the catalyst surface. From the experiments, information is also gained on the ratio between the rate of hydrogen adsorption and the rate of hydrogenation of methyl esters of fatty acids in vapor phase. Knowledge concerning the role of the catalyst in heterogeneously catalyzed reactions can be improved by fitting kinetic data to various rate equations. In order to calculate the action on the catalyst surface, it is necessary to assume a mechanism for the adsorption of the reactants. In spite of serious objections, the Langmuir-Hinshelwood adsorption isotherm is the most frequently applied mechanism for the adsorption process. The reaction between two compounds must be preceded by the adsorption of at least one of the reactants. It is possible to distinguish between three extreme cases, considering a chemical reaction where both of the reactants are adsorbed on the catalyst surface. 1. The reactants compete mutually for the active sites. 2. The reactants are adsorbed noncompetitively on the catalyst surface. 3. The presence of one adsorbed reactant enhances the adsorption of the other reactant. In case 1,according to Langmuir’s theory, the reactants are adsorbed on the same type of active sites and thus the adsorption is competitive. In the second case, the reactants are adsorbed on different sites on the surface. I t is also possible to explain the noncompetitive adsorption in the following way. If the difference in molecular weight is great between the two reactants, it is impossible for the larger molecules to cover the catalyst surface completely (Wells, 0888-5885/87/ 2626-0874$0l.50/0
1972). The third case, the enhanced adsorption, is explained by the formation of a complex between the two reactants on the catalyst surface (Weller, 1975). This surface complex is more strongly adsorbed than either component singly. Vapor-Phase Hydrogenation of Methyl Linolenate. Methyl linolenate is the methyl ester of a fatty acid containing 18 carbon atoms and 3 carbon-carbon double bonds. The vapor-phase hydrogenation of methyl linolenate, the triene, proceeds according to the reaction scheme (Magnusson, 1983) triene
I
I
-
H2
diene
HZ
+. monoene h
2%
I
-.
H2
saturate
(1)
The hydrogenations are catalyzed by supported copper and nickel catalysts. The objective of the present study was to elucidate how hydrogen and the methyl esters compete for the active sites on a-Alz03-supportednickel and copper catalysts. In applying the Hz/D2exchange reaction as a model reaction for the adsorption process, it was also possi.ble to calculate the ratio between the rate of adsorption for hydrogen and the rate of hydrogenation ac0 1987 American Chemical Society
Ind. Eng. Chem. Res., Vol. 26, No. 5, 1987 875 cording to reaction scheme 1. This provides very useful information considering the derivation of rate equations for the reaction system shown above. Experimental Section Experimental Procedure. Methyl linolenate was hydrogenated in vapor phase with a mixture of hydrogen and deuterium in a ratio of 1to 1 and in an excess of nitrogen. The hydrogenations were carried out in the presence of the supported copper and nickel catalysts, respectively. The kinetic study was performed in a gradient-free reactor described elsewhere (Lidefelt, 1983a). Each kinetic run consisted of two sequential periods. In the first period, the inlet flow to the reactor was the H2/D2 mixture in an excess of nitrogen. The conversion of the reaction Hz
+ Dz F? 2HD
(2)
was measured at the out-flow of the reactor. If reaction 2 is far from equilibrium, it is possible to calculate a net consumption rate for hydrogen (or deuterium). In the second period, the reactor in-flow consisted of methyl linolenate and of H2/D2,in the same concentration as in period 1. In this period, methyl linolenate was hydrogenated by the H2/D2 mixture, according to reaction scheme 1. The conversion of the hydrogenation reaction was measured at the out-flow of the reactor. Due to a large difference in concentration between the Hz/Dz mixture and methyl linolenate, the consumption of H2,Dz, and HD is small in the hydrogenation reactions. There are two reactions taking place on the catalyst surface, the Hz/Dz exchange reaction according to scheme 2 and the hydrogenation reaction according to scheme 1. The rate of the H2/D2exchange reaction in period 2 was thus measured in the presence of the methyl esters. Apparatus. The experimental equipment used in this investigation was the same as used in a previous study of Hz/D2 exchange as a model reaction for the hydrogen adsorption (Magnusson, 1987). The methyl ester was evaporated in a vaporization device described by Lidefelt (1983a). The concentration of the methyl ester was determined by the temperature in this unit. Analysis. The conversion of the H2/D2reaction was determined by using a quadrupole mass spectrometer of type Micromass 8 at the out-flow of the reactor. The analyses of the methyl esters were performed via a gas sampling valve, with a Perkin-Elmer F11 gas chromatograph using a Supelco 10% SP-2330 on 100/120 Chromosorb W 17 E 1-1851 column. Materials. The gases were supplied from gas cylinders and had a purity better than 99.98%. The methyl linolenate was of an analytical grade quality, better than 99.9% purity. The catalysts used were the same as those used in a previous study of hydrogen adsorption (Magnusson, 1987). These catalysts were also used in earlier investigations on vapor-phase hydrogenations of methyl esters of fatty acids (Magnusson, 1983; Lidefelt et al., 1983a-c). Computer Control. The entire experimental device was controlled by an ABC-80 microcomputer. The order of the runs was randomized by the computer. A more extended description of the experimental equipment is given elsewhere (Magnusson, 1983). Results and Discussion Twenty-four runs were performed in the presence of the nickel and copper catalyst, respectively. Each run consisted of the two sequential periods, described above. The experimental conditions are summarized in Table I. The
Table I. Experimental Conditions in the Nickel and Copper
Runs
temp range, " C H2/D2partial pressures in the reactor inlet, torr methyl linolenate partial pressure in the reactor inlet, torr w t of catalyst in reactor, g
catalyst Ni cu 148-188 170-216 19-88 64-277 0.1
0.44
2.29
6.75
Table 11. Observed and Predicted Decrease in H2/Dz Exchange Rate due to the Presence of the Methyl Esters on the Catalyst Surface. Nickel Catalyst pred" obsd rate of obsd decrease decrease in in exchange exchange H2/D2 exchange, rate, gmol s-l rate, rmol s-l temp. "C vmol ka-' ka-' ka-' 148
160
38.4 25.5 65.2 42.5 54.7 74.3 59.6 93.9 44.7 131.0 118.0 164.0
8.0 7.2
8.8 8.0 6.2 6.3 2.6 3.5 6.0 8.0 6.0 6.0
3.2 2.9 4.9 5.0 5.4 5.6 1.3 2.1 0.7 1.5 0.6 3.0
The predicted values are based on equilibrium constants from Lidefelt (1983b) and Magnusson (1983).
rate of hydrogenation was expressed as a total rate of the hydrogen consumption, according to reaction scheme 1. It was assumed that the rate of hydrogenation was equal for H2, Dz, and HD, respectively. The rate of the H2/D2 exchange reaction was expressed as the net consumption rate of hydrogen, according to eq 2. Decrease in the Rate of the Hz/Dz Exchange due to the Presence of the Methyl Esters on the Catalyst Surface. By analyzing the Hz/D2exchange reaction in the two sequential experimental periods, it was found that the presence of the methyl esters significantly decreased the rate of the H2/D2exchange reaction. In the presence of the nickel catalyst, the decrease in rate was 10-20% at low temperatures (148 "C) and 4-7% at high (188 "C). In the presence of the copper catalyst, however, the decrease in the exchange rate was only 2-14% and independent of the temperature. The interpretation of this observation was that the presence of the methyl esters on the catalyst surface decreases the number of active sites available for hydrogen adsorption. In a recent study of hydrogen adsorption (Magnusson, 1987), it was shown that the rate of the H2/Dzexchange reaction, in the presence of the nickel catalyst, followed the subsequent equation
In the presence of the copper catalyst, the rate of the exchange reaction was shown to follow the equation
The adsorbed methyl esters will, due to their large equilibrium constants (Lidefelt, 1983b; Magnusson, 1983), occupy a certain amount of the catalyst surface. If the methyl esters and hydrogen compete mutually for the active sites, the denominators, in eq 3 and 4, will contain
876 Ind. Eng. Chem. Res., Vol. 26, No. 5 , 1987 Table 111. Observed and Predicted Decrease in H2/D2 Exchange Rate due to the Presence of Methyl Esters on the Catalyst Surface. Copper Catalyst obsd decrease in predO decrease in predb decrease in obsd rate of H2/D, exchange rate, pmol exchange rate, kmol exchange rate, pmol temp, "C exchange, pmol s-l kg-* s-' kg-' s-l kg-' s-l kg-' 170 17.8 0.3 17.7 1.6 24.2 2.3 24.1 2.3 12.5 1.1 12.4 1.0 0 27.7 4.2 28.0 38.0 3.5 37.8 2.7 48.1 3.0 47.9 3.5 186 68.1 1.4 66.3 4.7 59.1 7.0 58.0 3.8 19.8 3.9 19.3 0.9 35.4 1.5 34.6 4.2 0 25.7 1.7 26.5 "The predicted values are based on equilibrium constants from Lidefelt (1983b) and Magnusson (1983). *The predicted values are based on an equilibrium constant fitted to t h e experimental data.
additional terms describing the adsorption of the methyl esters. The denominators in eq 3 and 4 will thus change to
Table IV. Ratio between the Rate of H2/D2Exchange and the Total Rate of Hydrogenation according to Reaction Scheme 1. Nickel Catalyst at P H p torr 57 62 10.4 7.0 21.7 16.6 25.8 21.1 37.2 30.0
rexchange/rhydrogenation
where p Land Ki are the partial pressures and equilibrium constants for all participating methyl esters. The experimental results were analyzed based on eq 3 for the nickel runs and eq 4 for the copper runs. The denominators in eq 3 and 4 were replaced by eq 5. The adsorption equilibrium constants for the methyl esters are known from previous adsorption studies (Lidefelt, 1983b; Magnusson, 1983). The results at two temperatures are presented for the nickel catalyst in Table 11. As may be seen in Table 11, the rate model eq 3 with eq 5 as the denominator predicts the decrease in exchange rate reasonably well. The conclusion must thus be that the methyl esters and hydrogen compete mutually for the active sites on the nickel catalyst. For the copper catalyst, the results are different. From the results presented in Table 111, it can be seen that eq 4, with eq 5 as the denominator and with known adsorption equilibrium constants, does not satisfactorily predict the decrease in exchange rate. Adsorption Equilibrium Constant for the Competitive Adsorption on the Copper Catalyst. In order to describe the influence which the presence of the methyl esters has on the exchange rate, an adsorption equilibrium constant, K,, was fitted to the experimental data. The constant was calculated to be K , = 111 f 18 atm-' (with 95% confidence limits). In Table 111, the calculated decrease in exchange rate, using K,, is presented. K,did not significantly depend on the temperature. The value of K , is to be taken as an average for the methyl esters present on the catalyst surface. K , is thus the adsorption equilibrium constant for the methyl esters adsorbed on sites for which the methyl esters and hydrogen compete mutually. The value of K , is to be compared with the reported adsorption equilibrium constants on copper in the absence of hydrogen (Lidefelt, 198313). For methyl linoleate, the values are 25.3 X lo3 and 9.7 X los atm-' at 170 and 186 "C,respectively. A possible interpretation of the relative insensitivity of the H2/D2exchange rate to the presence of methyl esters is that hydrogen and the methyl esters only compete for a minor part of the active sites on the copper catalyst. It would also appear that the competitive adsorption of the methyl esters has a low enthalpy of adsorption, since K , did not significantly depend upon the temperature. Comparison between the Hz/D2Exchange and the Rate of Hydrogenation of Methyl Linolenate. In deriving rate equations for the reactions in eq 1,it is of great
temp, "C 148 161 175 188
19 3.3 6.2 7.9 10.3
31 7.0 11.3 16.3 20.2
39 5.0 12.7 16.4 23.0
88 12.3 27.4 34.7 47.6
Table V. Ratio between the Rate of H2/D2Exchange and the Rate of Hydrogenation according to Reaction Scheme 1. Copper Catalyst at PHp torr 166 209 10.3 17.3 14.5 19.6 35.2 35.9 53.3
rerchange/rhydrogenation
temp, "C 170 186 202 216
64 5.2 5.4 7.1 11.4
119 7.0 10.2 15.0 23.1
121 11.1 15.7 20.9 34.3
277 21.6 31.9 48.3 76.0
interest to know how fast the adsorption of hydrogen is, compared with the rate of hydrogen consumption in the hydrogenation reactions. In Tables IV and V, the ratio between the rate of hydrogenation is shown for various hydrogen pressures and temperatures. The hydrogen pressures in the tables are the sum of the partial pressures for Hz, Dz, and HD. The rate of the exchange reaction is coupled to the rate of hydrogen adsorption. For copper, eq 4, the exchange rate equals the rate of adsorption for hydrogen (Magnusson, 1987). The total rate of adsorption for hydrogen plus deuterium is thus twice the exchange rate. For the nickel catalyst, eq 3, the relation between the rate of adsorption and the exchange rate seems somewhat more disputable, cf. Magnusson (1987). It can be concluded, however, that twice the exchange rate should define an upper limit for the adsorption rate. From the results in Tables IV and V, it can be seen that the activation energy for the exchange reaction is higher than for the hydrogenation reaction, since the relative rate of the exchange to the hydrogenation reaction increases with increasing temperature for both catalysts. It may also be noted that the reaction order with respect to hydrogen is higher for the exchange reaction. Although the experimental conditions were different with respect to temperature and hydrogen pressure in the copper and nickel runs, there are resemblances between the results in Tables IV and V. It must be emphasized, however, that under equal conditions, the rate of the Hz/D2 exchange in the presence of the nickel catalyst exceeds the rate in the presence of the copper catalyst by a factor of 20. Conclusion By analyzing the rate of the H2/D2exchange with and
Ind. Eng. Chem. Res. 1987,26, 877-881
877
without methyl linolenate present, it was possible to calculate the competition for the active sites on the catalysts studied. It was found that hydrogen and the methyl esters compete mutually for the active sites on the nickel catalyst. Based upon the calculated value of the adsorption equilibrium constant for the methyl esters, it was concluded that hydrogen and the methyl esters only compete for a minor part of the active sites on the copper catalyst. In earlier investigations (Magnusson, 1987) it was concluded that the adsorbed state is more mobile on the copper catalyst than on the nickel catalyst. This was valid for hydrogen, as well as for the methyl esters. The higher mobility on the copper catalyst could be a possible explanation to the noncompetitive adsorption on the copper catalyst. The comparison between the rate of Hz/D2exchange and the rate of hydrogenation, according to reaction scheme 1, indicates that the commonly applied assumption of equilibrium adsorption, in deriving rate equations, could be disputable. At low temperatures and low hydrogen pressures, the ratio between the rate of adsorption for hydrogen and the rate of hydrogenation is merely a factor of 10. Moreover, the difference in activation energy and hydrogen dependence for the two processes indicates a possible shift in the rate-controlling step for the hydrogenation of methyl linolenate. Acknowledgment
Nomenclature k = rate constant in eq 3, mol s-l kg-' atm-' k , = adsorption rate constant in eq 4,mol kg-' atm-' K, = adsorption equilibrium constant methyl esters on copper catalyst, atm-' (K, is associated with the competitive adsorption) KHz = adsorption equilibrium constant for hydrogen, atm-' pHz = partial pressure of hydrogen, atm-' p H D = partial pressure of HD, atm-' rHz= net consumption rate of hydrogen in the exchange reaction, mol s-' kg-'
The Swedish Board for Technical Development has provided financial support.
Received for review February 27, 1986 Accepted November 17, 1986
Registry No. Methyl linoleate, 301-00-8; Cu, 7440-50-8; Ni, 7440-02-0;Hz,1333-74-0.
Literature Cited Lidefelt, J.-0. J. Am. Oil Chem. SOC. 1983a, 60, 588. Lidefelt, J.-0. J. Am. Oil Chem. SOC. 1983b, 60, 593. Lidefelt, J.-0.;Magnusson, J.; Schoon N.-H. J.Am. Oil Chem. SOC. 1983a, 60, 600. Lidefelt, J.-0.;Magnusson, J.; Schoon N.-H. J. Am. Oil Chem. SOC. 1983b, 60, 603. Lidefelt, J.-0.;Magnusson, J.; Schoon N.-H. J.Am. Oil Chem. SOC.
19830,60,608.
Magnusson, J. Doctoral thesis, Chalmers Univeristy of Technology, Goteborg, Sweden, 1983. Magnusson, J. Ind. Eng. Chem. Res. 1987, in press. Weller, S. W. Adu. Chem. Ser. 1975, 148, 1. Wells, P. B. In Surface and Defect Properties of Solids; The Chemical Society: London, 1972; Vol. 1.
H2/D2 Exchange as a Model Reaction for Studying Hydrogen Adsorption on a-A1203-SupportedCopper and Nickel Catalysts Jeppe Magnusson Department of Chemical Reaction Engineering, Chalmers University of Technology, S-412 96 Goteborg, Sweden
A method of measuring adsorption enthalpy and entropy for hydrogen on a-A1203-supportedcopper and nickel catalysts is described. T h e H2/Dz exchange reaction, catalyzed by these catalysts, was shown to be a model system for the hydrogen adsorption process. The results obtained were compared with equilibrium chemisorption measurements of hydrogen separately performed in static vacuum equipment. Information is also given about the kinetics of adsorption. Introduction Scope. Kinetic measurements are a powerful tool for the understanding of the mechanisms of heterogeneously catalyzed reactions. Kinetic models, based upon the Langmuir-Hinshelwood adsorption isotherm, are often fitted to the experimental data in order to explain the action on the catalysts surface. These so-called mechanistic models contain adsorption parameters and parameters associated with surface reactions. If it is possible to determine the adsorption parameters with independent physical methods, only the rate constants for the surface reactions remain to be estimated from kinetic experiments. A nonbiased fitting of kinetic data is thus a strong support for a proposed mechanistic model. The hydrogenation of vegetable oil plays an important role in the production of edible fats. In a recent study (Lidefelt et al., 1983a-c) of vapor-phase hydrogenation of methyl oleate and methyl linoleate over supported nickel
and copper catalysts, the selective actions of the catalysts were elucidated. It was shown that the ability of the catalysts to adsorb hydrogen determined the difference in selectivity between the catalysts studied. The objective of this paper is to present a new technique to measure adsorption quantities for hydrogen on supported metal catalysts. Adsorption Studies of Hydrogen. Over the last 5-6 decades, many investigations have been reported on the adsorption of hydrogen on nickel and copper. Results obtained prior to x 5e early 1960s are summarized by Bond (1962) and Trapnell and Hayward (1964). A more recent review is given by Toyoshima and Somorjai (1979). Much attention has been given to the enthalpy of adsorption and less to the entropy, as well as to the rate and the activation energy of adsorption. Moreover, most studies have been performed under extreme conditions, such as very low pressure and temperature. Due to the great difference in
0888-5885/87/2626-0877$01.50/0 0 1987 American Chemical Society
