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Ind. Eng. Chem. Res. 1997, 36, 2970-2975
Dynamics of Catalytic Methane Coupling† Y. Mortazavi,*,‡ R. R. Hudgins, and P. L. Silveston* Department of Chemical Engineering, University of Waterloo, Waterloo, Ontario, Canada N2L 3G1
Transient experiments at 750 °C and 1 atm with methane and O2 over Li2O/MgO and CeO/ Li2O/MgO catalysts demonstrate that coupling can occur in the absence of gas-phase oxygen. Cerium oxides increase the amount of oxygen available for the reaction. CO2 also forms from oxygen associated with the catalyst and appears to react on the surface to produce a carbonate. Oxygen in the gas phase is needed to decompose this carbonate at 750 °C. Adsorbed oxygen leads mainly to methane combustion, while most of the oxygen from the catalyst structure results in C2 products. Introduction Although methane coupling has lost the glamour it enjoyed for much of the last decade, research questions remain, the answers to which could be useful when this topic reappears in the expected cycle of research fashions. Our main interest in this paper is the transient behavior of methane coupling when a step change in a reactant concentration is introduced. Among others, Kobayashi and Kobayashi (1974), Bennett (1976), and Kobayashi (1982) have shown that dynamic measurements of this type provide useful insight into chemical mechanisms. Such measurements often can identify rate-determining steps. In this paper, we apply dynamic measurements to examine mechanisms proposed by others. Two widely investigated catalysts, Li2O/MgO and CeO/Li2O/MgO, were used (Ito et al., 1985, 1987; Driscoll et al., 1985; Moriyama et al., 1986; Mirodatos et al., 1987; Hutchings et al., 1987, 1989a,b; Campbell and Lunsford, 1988; Lin et al., 1988; Cant et al., 1988; Roos et al., 1989; Nelson et al., 1989; Hargreaves et al., 1991). Experimental System All experiments were performed at 750 °C and slightly above atmospheric pressure in the quartz microreactor shown in Figure 1. After the catalyst bed, the reactor narrows and has been stuffed with quartz wool to reduce the volume available for homogeneous reactions. The microreactor was mounted in a tube furnace, the temperature of which could be controlled to (2 °C. For the step-change experiments, 0.40 g of catalyst was used. The quartz-wool-filled tube was extended out of the furnace, so cooling was rapid, and was connected to either a 16-port sampling valve or a 10-port sampling valve with 2 sampling loops. The second system allowed injection into gas chromatographs in parallel. One of these employed a Porapak N column and a thermal conductivity detector (TCD) and was used to analyze for O2 and CO2. Following the TCD, the stream passed to a flame ionization detector (FID) to determine the hydrocarbons in the sampling loop. The second gas chromatograph, used to measure CO, was equipped with a TCD and a column packed with 5-Å molecular sieves. Reactor temperatures and pressure were also measured. † This contribution is dedicated to Dr. Gilbert Froment, a pioneer of chemical reaction engineering. The paper is submitted in recognition of his contributions to the use of dynamic measurements to elucidate mechanisms and rate-determining steps. ‡ Present address: Chemical Engineering Department, University of Tehran, Tehran, Iran.
S0888-5885(96)00604-5 CCC: $14.00
Figure 1. Schematic diagram of the microreactor used for the experiments.
A constant total flow rate of 200 mL/min (STP) was used. All transient experiments were begun after steady state was attained at 750 °C, using a mixture of 10% methane and 5% O2, with helium making up the balance. Step changes were introduced by three-way solenoid valves with carefully balanced lines to avoid pressure surges on switching. Sample collection was done at 4-s intervals and was initiated 30 s after a composition switch. This allowed for transport lag as well as mixing. There has been a controversy in the literature concerning the mechanism of methyl radical formation. Driscoll et al. (1985) argue that the radical is formed by the interaction of methane with a surface, possibly with an adsorbed oxygen ion radical O-, whereas, Sokolovskii et al. (1989) assume that activation occurs heterolytically via an acid-base group involving lattice oxygen. To investigate catalyst participation in methyl radical formation, we chose two catalysts, Li2O/MgO and CeO/Li2O/MgO. The former is generally assumed irreducible, while Ce in the second catalyst provides an oxygen reservoir. © 1997 American Chemical Society
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then decline. The initial selectivity to C2H6 is good and much higher than the selectivity to C2H4. C2H4 vanishes from the off-gas after about 20 s, but ethane persists for over 1 min. CO2 is produced at a rate of about 0.2 mmol/(gcat‚h) 60 s after the switch to methane, even though there is no O2 in the feed. In Figure 2b, the response of the Ce-containing catalyst to the introduction of methane is shown. The initial peaks for ethane and ethene are higher with this catalyst, and these products persist for much longer times after the introduction of methane. Although the peak for CO2 is about the same for both catalysts, CO2 production continues at a higher rate for the CeO/Li2O/MgO catalyst. It is evident from Figure 2a that methane reacts with oxygen within the Li2O/MgO catalyst, producing C2H4, C2H6, and CO2. The C2H6/C2H4 ratio at the time of the switch was about 14.6 and increased with time after methane was introduced. These observations suggest that ethene may not be a primary coupling product over Li2O/MgO. Furthermore, the low ratio suggests that dehydrogenation of ethane, thought to occur in the gas phase, is suppressed by the presence of the catalyst. It appears that the role of nonadsorbed surface oxygen in different catalysts is not the same. For instance, lithium-doped TiO2 and Sm2O3 were found to activate methane in the absence of gas-phase oxygen but led solely to the production of CO2 (Efstathiou et al., 1993; Ekstrom and Lapszewicz, 1989). On the other hand, Pr6O11 produced both C2H6 and CO2 (Ekstrom and Lapszewicz, 1989). Ito et al. (1985) propose the following redox cycle for ethane formation:
Figure 2. Response to step change from 0 to 10 vol % methane in the absence of O2 in the feed for an oxidized catalyst previously flushed by He. Operating conditions: 750 °C, 1.0 atm, 30 L/(gcat‚h) (STP).
Both catalysts were prepared from published recipes (Ito et al., 1985). Before use, catalyst samples were stabilized at 750 °C by exposure to a mixture of CH4: O2:He in the proportions 20:10:170 mL/min. Direct coupled plasma spectrometry (DCP), X-ray diffractometry (XRD), and X-ray photoelectron spectroscopy (XPS) were used to characterize the catalyst and to monitor bulk and surface composition changes through calcining and extended use under reaction conditions. Results and Discussion To investigate the participation of surface or lattice oxygen, the catalyst was switched from steady state as described above to 5% O2 in He for 30 min. After just 1 or 2 min, the off-gas contained no carbon oxides or hydrocarbons. Then the catalyst was exposed to 200 mL/min (STP) of pure helium for another half-hour. Finally, a switch to 10% CH4/He was made. Researchgrade methane with a purity of 99.99% was used, and an oxygen trap was installed upstream from the methane mass flow controller. The response for this step change is shown in Figure 2. Immediately after the switch to 10% CH4, CO2, ethane, and ethene concentrations reach maxima and
O- + CH4 f OH- + CH3•
(1)
CH3• + CH3• + M f C2H6
(2)
2OH- f O2- + 0 + H2O
(3)
O2- + 0 + 1/2O2 f 2O-
(4)
where 0 represents an anion vacancy. Peil et al. (1991) investigated oxidative coupling over Li2O/MgO by means of steady-state isotopic transients using 18O2 and 13CH4. They concluded that the catalyst provides parallel pathways for the conversion of methane. One pathway is active in the formation of carbon oxides, while the second is active in the formation of C2H6. Furthermore, they found that up to 12 subsurface layers of the catalyst participate in the reaction. The response for CeO/Li2O/MgO is similar to that of the Li2O/MgO catalyst except for higher rates of C2H4 and C2H6 formation in the first few seconds after introducing CH4 and the persistence of C2H4 formation. The decay of all product formation rates is slower than for the Li2O/MgO catalyst. Thus, CeO/Li2O/MgO catalyst is capable of providing more oxygen for the coupling reaction per unit mass of catalyst than Li2O/MgO. The extent of the bulk-phase oxygen contribution is significantly increased by addition of CeO2 to Li2O/MgO. Because cerium has two oxidation states, III and IV, the redox cycle probably involves these states. Although the step-change experiments were not designed to determine the amount of oxygen available in the catalysts for methane coupling, the areas under the curve in Figure 2 indicate that 0.02 mmol of atomic oxygen is removed from each gram of Li2O/MgO catalyst by the surface reactions within the first 60 s after the
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Figure 3. Step-up and step-down experiments: (a, b) step up of O2 from 0 to 5 vol % in the presence of 10 vol % methane in the feed; (c, d) step down of methane from 10 to 0 vol % after steady-state operation with 10 vol % methane and 5 vol % O2; (e, f) step down of O2 from 5 to 0 vol % after steady-state operation with 10 vol % methane and 5 vol % O2. (a, c, e) For Li2O/MgO catalyst; (b, d, f) for CeO/ Li2O/MgO catalyst. Operating conditions as in Figure 2.
step change. For the CeO/Li2O/MgO catalyst, the amount of oxygen is 0.026 mmol/g. These amounts are very much less than the oxygen available in the catalyst, which we estimate from the composition measurements to be 31 and 18 mmol, respectively. From this comparison, it is clear that the Ce promoter contributes a substantially higher fraction of the oxygen available in the catalyst to the coupling reaction. Clearly, combustion of methane or oxidation of C2 products proceeds with oxygen from the catalyst. Either or both of these reactions are the only source of CO2 found in the off-gas. A second set of experiments used a step up in oxygen from 0 to 5 vol % in the presence of 10 vol % methane. Prior to this step change, the catalyst was exposed to 10% CH4/He gas for 30 min at a total flow rate of 200 mL/min (STP). No carbon oxides were detected in the off-gas before the introduction of O2, so it was assumed that the catalyst was fully reduced. Then the catalyst was flushed with pure helium at 200 mL/min (STP) for another 30 min. After this, there was a 10-min exposure of 10% CH4/He; finally, the switch to a 10% CH4-5% O2/He mixture was made. The first two panels of Figure 3 (i.e., parts a and b) depict the responses of the Li2O/MgO and CeO/Li2O/ MgO catalysts respectively to the step change just described. C2H4 and C2H6 appear instantly after O2 is introduced, and their rates of formation reach steady state within seconds. CO2 also appears instantly, but its rate of formation gradually increases before leveling off after almost 40 s. The CO2 observation suggests that CO2, just as C2H4 and C2H6, is formed initially at a rate close to its steady-state rate but some of the CO2 is adsorbed by the catalyst to form a surface carbonate. Our XPS measurements appear to support this expla-
Table 1. XPS Surface Composition (Atomic Mass %) of the Catalyst Samples elements detected sample
Mg
O
C
fresh Li/MgO used Li/MgO fresh Ce/Li/MgO used Ce/Li/MgO fresh Li/MgO (after 2-min sputtering) fresh Ce/Li/MgO (after 2-min sputtering)
54.0 47.0 48.0 37.0 61.0
32.0 28.0 29.0 28.0 30.0
14.0 25.0 22.0 34.0 8.7
58.0
33.0
7.8
Ce
Li
1.0 1.0 1.0
nation. Elements detected on the surface of fresh and used catalysts are given in Table 1. It can be seen that the carbon concentration on the surface of the used catalyst is substantially higher than that of the fresh one. XRD observations showed that lithium in both catalysts was present partially in the form of Li2CO3 (Mortazavi, 1994). Literature evidence further supports our explanation for the increase in the CO2 rate of formation with time after the switch. Temperature-programmed desorption (TPD) studies of the CeO/Li2O/MgO catalyst have shown that CO2 adsorbs strongly on the surface and does not desorb completely below 800 °C (Bartsch and Hofmann, 1990). Peng et al. (1990), using the same recipe for Li2O/MgO catalyst preparation as used in our work, observed no carbonate carbon for fresh catalyst. However, further treatment under reaction conditions, i.e., exposing the catalyst to CH4/O2 at 650 °C, led to the appearance of a carbonate in the XPS spectra that was attributed to lithium carbonate formation. The remainder of Figure 3 shows the response to a step down in each of the reactants after steady state at 10% CH4/5% O2 was reached. Thus, parts c and d of
Ind. Eng. Chem. Res., Vol. 36, No. 8, 1997 2973
Figure 3 show the response to a step down in methane from 10 to 0 vol % in the presence of 5 vol % oxygen, while parts e and f of Figure 3 give the response to a step down in oxygen from 5 to 0 vol % in the presence of 10 vol % methane. These two panels are the counterparts of parts a and b of Figure 3. Although the step-down experiments seem to show similar responses, there are differences depending on which reactant is withdrawn from the feed gas. When methane is removed from the feed, C2 production disappears within 5 s. CO2, on the other hand, persists for as much as 50 s (the duration of our observations), indicating either strong adsorption of CO2, probably in the form of carbonate, or deposition of carbon on the catalyst surface. The amounts are slightly higher for the unpromoted Li2O/MgO catalyst initially but decay to virtually identical rates of formation at 50 s. After methane disappears from the feed, only small amounts of C2H4 and C2H6 are detected and only during the first few seconds. Consequently, neither C2H4 nor C2H6 adsorb appreciably on the surface, or if adsorbed, they are oxidized immediately to CO2. For O2 step down (see Figure 3e and 3f), the CO2 response is similar to that of the step down of methane shown in the previous two panels. The rate of decay, however, is slower so that the CO2 concentration after 50 s is about 50% greater for the O2 step down. The presence of CO2 in the off-gas in the absence of O2 in the feed indicates that CO2 is not coming from the combustion of carbon residues on the surface. CO2 must result either from desorption or from CH4 surface combustion or further oxidation of the C2 products by lattice oxygen. Ethane continues to form for both catalysts as much as 60 s after the O2 step down. This observation confirms that the catalysts are able to provide oxygen for dehydrogenation and formation of methyl radicals. The amount of C2H6 formed immediately after the step down is about the same for both catalysts, but the decay in C2H6 formation is slower for the Ce-promoted material. After 50 s, the rate of formation of C2H6 is 3-fold the rate observed for the Li2O/MgO catalyst. Unlike Figure 2, C2H4 vanishes after 5 s for both catalysts, suggesting that gas-phase O2 is more important for C2H4 formation than oxygen from the catalyst surface. An alternative explanation is that C2H4 formation involves sites of greater activity that are reduced quickly after oxygen is removed from the feed stream. Differences in the response between Figure 2 and Figure 3e and 3f, even though both represent the removal of O2 from the gas phase, suggest that two forms of oxygen (adsorbed and bulk) may be active in methane coupling. To investigate the role of each of these oxygen sources, the methane step-up experiment in the absence of gasphase O2 was repeated but without stripping off adsorbed oxygen with a 10-min He flush of the system. The response (shown as solid symbols) is compared to that with stripping (shown as open symbols) in Figure 4. The much higher formation rates for CO2, when stripping is not used, means that adsorbed oxygen is participating in the methane coupling reaction. However, the C2H6 formation rates show a small increase for the Li2O/MgO catalyst and appear to be unchanged for the Ce catalyst. With respect to C2H4, there appears to be no effect of the He flush. Consequently, adsorbed oxygen must combust methane although it is conceivable that it has a dual role of forming the methyl radical and destroying these radicals so that the net effect on
Figure 4. Comparison of responses for a step up from 0 to 10 vol % methane for an oxidized catalyst in the absence of O2 in the feed with and without flushing of catalyst surface. Operating conditions as in Figure 2.
C2 formation is very small. The persistence of adsorbed oxygen on the surface for at least 1 min is demonstrated by the failure of the CO2 curve to drop to the level of the CO2 curve when flushing is used. Eventually, both response curves must fall to zero as oxygen associated with the catalyst is depleted. Independent measurements of lags and mixing times in the experimental system established that the responses cannot be due to mixing effects. Certainly Figure 4 indicates that the adsorbed oxygen and lattice oxygen-forming part of the catalyst surface behave differently. The latter leads primarily to the methyl radical, which desorbs and forms ethane in a gas-phase reaction. Ito et al. (1985) propose that a methyl peroxy radical can be formed on the surface in the presence of O2-CH4 mixtures, leading to CO and CO2 on desorption. We suggest that adsorbed oxygen leads primarily to the methyl peroxy radical. Curiously, the very low formation rates of C2H4 indicate that neither adsorbed O2 nor oxygen from the catalyst bulk participates in C2H4 formation. In the literature, there is evidence that the role of adsorbed oxygen depends on the type of catalyst. Keulks and Yu (1987), working with a potassium-promoted bismuth phosphate catalyst,
2974 Ind. Eng. Chem. Res., Vol. 36, No. 8, 1997
Figure 5. Response to a discontinuation of reactant flow from steady state for CeO/Li2O/MgO catalyst. Operating conditions as in Figure 2.
observed effects similar to those shown in Figure 4. Imai et al. (1987) used TPD on a LaAlO3 catalyst to conclude that a strongly adsorbed oxygen was involved in methane coupling, whereas weakly adsorbed oxygen led to total combustion. Although TPD measurements were not made in this study, step-change experiments provide similar information. In this case, they indicate that adsorbed and bulk-phase oxygen must be active in methane coupling under the steady-state operation used by most investigators of the Li2O/MgO catalysts. Integrating under the curves in Figure 4 indicates that 0.05 mmol of atomic oxygen is removed per gram of the Li2O/ MgO catalyst, while 0.054 mmol is removed per gram of promoted catalyst in the first 60 s after oxygen is removed from the feed gas. This is almost twice the oxygen removed from the catalyst when flushing is carried out. The extent of CH4 adsorption on one of the catalysts was investigated by a step-change experiment employing the CeO/Li2O/MgO catalyst in which both reactants, CH4 and oxygen, were abruptly removed from the feed stream. The response for this step change appears in Figure 5. Figure 5a shows that CH4 disappears essentially instantaneously from the reactor off-gases. The very small amounts of CH4 detected probably arise from CH4 adsorption elsewhere in the reaction system. Figure 5b shows that C2H6 and C2H4 vanish within seconds of the step change. The surprising observation in Figure 5b is that the CO2 in the off-gas is very small. Indeed, areas under the response curves show just about 0.01 mm of CO2/g of catalyst swept out in the first 60 s after a step down in O2. Thus, in the absence of oxygen, the Li2CO3 that appears from our XPS measurements to be present on the catalyst does not decompose. On the other hand, Figures 2 and 3 show that when O2 is present in the gas phase, decomposition must occur. Our conclusion that O2 must be present for Li2CO3 decomposition is supported by the observation that the activation of the Li2O/MgO catalyst with lithium carbonate as the Li precursor is accomplished by calcining the catalyst under oxygen (Ito et al., 1985). XPS studies of Peng et al. (1990) found that the surface Li was not in the form of lithium carbonate after calcination under an oxygen atmosphere. Perhaps our experimental temperature of 750 °C was too low to decompose lithium carbonate in the absence of O2.
Acknowledgment Support for this project came from operating grants from the Natural Sciences and Engineering Research Council of Canada to two of us (R.R.H., P.L.S.). We also acknowledge partial support through a fellowship from the Government of Iran (to Y.M.). Literature Cited Bartsch, S.; Hofmann, H. Investigations on a Ce/Li/MgO-Catalyst for the oxidative coupling of methane. Catal. Today 1990, 6, 527. Bennett, C. O. The transient method and elementary steps in heterogeneous catalysis. Catal. Rev.sSci. Eng. 1976, 13, 121. Campbell, K. D.; Lunsford, J. H. Contribution of gas-phase radical coupling in the catalytic oxidation of methane. J. Phys. Chem. 1988, 92, 5792. Cant, N. W.; Lukey, C. A.; Nelson, P. F.; Tyler, R. J. The rate controlling step in the oxidative coupling of methane over a lithium-promoted magnesium oxide catalyst. J. Chem. Soc., Chem. Commun. 1988, 766. Driscoll, D. J.; Martir, W.; Wang, J.-X.; Lunsford, J. H. Formation of gas-phase methyl radicals over MgO. J. Am. Chem. Soc. 1985, 107, 58. Efstathiou, A. M.; Papageorgiou, D.; Verykios, X. E. Transient kinetic study of the reaction of CH4 and C2H6 with the lattice oxygen of Li+-doped TiO2 catalyst. J. Catal. 1993, 141, 612. Ekstrom, A.; Lapszewicz, J. A. A study of the mechanism of the partial oxidation of methane over rare earth oxide catalysts using isotope transient techniques. J. Phys. Chem. 1989, 93, 5230. Hargreaves, S. J.; Hutchings, G. J.; Joyner, R. W. Hydrogen production in methane coupling over magnesium oxide. In Natural Gas Conversion; Holmen, A., Jens, K.-J., Kolboe, S., Eds.; Elsevier: Amsterdam, 1991; p 155. Hutchings, G. J.; Scurrell, M. S.; Woodhouse, J. R. The role of surface O- in the selective oxidation of methane. J. Chem. Soc., Chem. Commun. 1987, 1388. Hutchings, G. J.; Scurrell, M. S.; Woodhouse, J. R. Partial oxidation of methane over samarium and lanthanum oxides: a study of the reaction mechanism. Catal. Today 1989a, 4, 371. Hutchings, G. J.; Scurrell, M. S.; Woodhouse, J. R. The role of surface O- in the selective oxidation of methane. Chem. Soc. Rev. 1989b, 18, 251. Imai, H.; Tagawa, T.; Kamide, N. Oxidative coupling of methane over amorphous lanthanum aluminum oxides. J. Catal. 1987, 106, 394. Ito, T.; Wang, J.-X.; Lin, C.-H.; Lunsford, J. H. Oxidative dimerization of methane over a lithium-promoted magnesium oxide catalyst. J. Am. Chem. Soc. 1985, 107, 5062. Ito, T.; Tashiro, R.; Watanabe, T.; Toi, K.; Ikemoto, T. Activation of methane on the MgO surface at low temperatures. Chem. Lett. 1987, 1723. Keulks, G. W.; Yu, M. The oxidative coupling of methane. React. Kinet. Catal. Lett. 1987, 35, 361.
Ind. Eng. Chem. Res., Vol. 36, No. 8, 1997 2975 Kobayashi, M. Characterization of transient response curves in heterogeneous catalysissI. Classification of the curves. Chem. Eng. Sci. 1982, 37, 393. Kobayashi, H.; Kobayashi, M. Transient Method in Heterogeneous Catalysis. Cat. Rev.sSci. Eng. 1974, 10, 139. Lin, C.-H.; Wang, J.-X.; Lunsford, J. H. Oxidative dimerization of methane over sodium-promoted calcium oxide. J. Catal. 1988, 111, 302. Mirodatos, C.; Perrichon, V.; Durupty, M. C.; Moral, P. Deactivation of alkali promoted magnesia in oxidative coupling of methane. In Catalyst Deactivation; Delmon, B., Froment, G. F., Eds.; Elsevier: Amsterdam, 1987; Vol. 34. Moriyama, T.; Takasaki, N.; Iwamatsu, E.; Aika, K. Oxidative dimerization of methane over promoted magnesium oxide catalysts. Chem. Lett. 1986, 1165. Mortazavi, Y. Steady-state transient and periodic operation studies of oxidative coupling of methane. Ph.D. Thesis, University of Waterloo, Waterloo, Ontario, Canada, 1994. Nelson, P. F.; Lukey, C. A.; Cant, N. W. Measurements of kinetic isotope effects and hydrogen/deuterium distributions over methane oxidative coupling catalysts. J. Catal. 1989, 120, 216. Peil, K. P.; Goodwin, J. G., Jr.; Marcelin, G. Surface phenomena during the oxidative coupling of methane. J. Catal. 1991, 131, 143.
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Received for review October 1, 1996 Revised manuscript received January 17, 1997 Accepted January 24, 1997X IE9606045
X Abstract published in Advance ACS Abstracts, June 15, 1997.
