Kinetics and Hydrogen Removal Effect for Methanol Decomposition

488

Ind. Eng. Chem. Res. 1999, 38, 488-492

SEPARATIONS Kinetics and Hydrogen Removal Effect for Methanol Decomposition S. Hara,* W.-C. Xu, K. Sakaki, and N. Itoh National Institute of Materials and Chemical Research, Tsukuba 305-8565, Japan

Methanol decomposition to carbon monoxide in a palladium membrane reactor (PMR) is presented, where a 1 wt % Pd/SiO2 catalyst and a Pd91Ru6In3 alloy membrane tube were employed. Experiments were carried out at relatively low temperatures, 220-250 °C, in a PMR mode as well as a conventional catalytic reactor (CCR) mode. A kinetic analysis for methanol conversion change in the CCR mode revealed that the overall reaction rate was controlled by the desorption process of the resultant carbon monoxide from the active sites of the catalyst surface. In the PMR mode it was shown that the selective separation of produced hydrogen from the reaction to the permeate side led to an increase in methanol conversion. Further, amounts of byproducts such as carbon dioxide, dimethyl ether, and methyl formate were found to be also influenced by the hydrogen separation. 1. Introduction Membrane reactors, containing a catalyst and a permselective membrane in one reactor, have been successfully examined for many reactions not only to produce useful chemicals for industries such as benzene,1 ethylene,2 and propylene,3 but also to decompose harmful compounds for the environment.4 Moreover, these are largely expected to be applied to dehydrogenation of chemical hydrogen carriers such as cyclohexane, methylcyclohexane, etc., which is an important step in the global hydrogen transportation cycles.5,6 Although the dehydrogenations of these chemicals are endothermic and equilibrium-limited reactions, requiring rather high operating temperatures, it has already been demonstrated that the membrane reactors could release them from the limitation and achieve a high conversion even at relatively low temperatures.7-9 Methanol is also considered as a promising candidate for the chemical carrier to transport not only hydrogen but also heat, because it can easily be produced from and decompose to carbon monoxide and hydrogen with a relatively large enthalpy change (90.7 kJ/mol). Nevertheless, the methanol decomposition also favors high temperatures. It is, therefore, strongly required to proceed smoothly at low temperatures for effective use of waste heat sources. In this study, the methanol decomposition to carbon monoxide and hydrogen was carried out in a palladium membrane reactor (PMR) in order to promote the reaction at low temperatures, where a dense hydrogenpermeable membrane of palladium alloy and a Pd/SiO2 catalyst were employed. First, the permeation behavior of the membrane and the decomposition kinetics over the catalyst were analyzed in detail. Next, the reaction behavior in the PMR was compared with those in a conventional catalytic reactor (CCR). Finally, the influence of hydrogen removal on selectivities to a small amount of byproducts was also discussed. * To whom correspondence should be addressed. Tel.: +81298-54-4663. Fax: +81-298-54-4674. E-mail: [email protected].

Figure 1. Schematic of a double-tube membrane reactor including a hydrogen-permeable membrane of Pd alloy.

2. Experimental Section Figure 1 is the double-tube membrane reactor employed in this study, which was in a thermostat to keep the temperature constant. This reactor consisted of a stainless steel shell tube (22 mm i.d.) and a hydrogenpermeable membrane tube made of Pd91Ru6In3 alloy (0.2 mm thickness, 7.1 mm o.d., 208 mm length). A 1/16-in. pipe was inserted inside the membrane tube to exhaust sweep gas containing hydrogen permeated through the membrane. In the annular space around the membrane, the catalyst pellets (1 wt % Pd/SiO2, 1.1-1.6 mm), supplied by Kawasaki Heavy Industries, Ltd., Japan, were packed. A gaseous mixture of argon and methanol, which was fed by a microsyringe in the range of 0.8838 µmol/s and evaporated in the thermostat, was introduced over the catalyst layer. The gaseous components in the effluent were analyzed by a gas chromatograph (Shimadzu GC-14A) with an activated carbon column and a thermal conductivity detector mainly for inorganic gases and a Porapack T column and a flame ionization detector for organic compounds. The decomposition was performed for three different modes: (i) closing the inlet and the outlet of the permeate side to prevent hydrogen permeation (CCR mode), (ii) feeding Ar gas at a rate of 10 cm3/min into the permeate side cocurrently to sweep the permeated hydrogen (PMR sweep mode), (iii) evacuating the inside

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of the membrane by a rotary vacuum pump, thereby keeping a permeate-side pressure of less than 70 Pa at the outlet (PMR vacuum mode). It took more than 2 h before a steady result was obtained. In particular, the reaction in the vacuum mode at elevated pressures was often unstable, so that several hours was necessary to obtain the reliable data. In the CCR mode, the influences of methanol concentration in the feed and feed rate on methanol conversion were examined to formulate the reaction rate. The effect of the hydrogen removal from the reaction region on the decomposition was investigated at 220 °C using 100% methanol as the feed. To characterize the hydrogen-permeable membrane, pure hydrogen was induced into the reaction side at pressures of 0.12-0.2 MPa and the hydrogen flow rate from the outlet of the permeate side was measured by a soap flowmeter. No Ar permeation was detected by this method, indicating that the membrane was pinholefree. 3. Results and Discussion 3.1. Hydrogen Permeation Properties. The hydrogen permeation rate of the Pd alloy membrane employed in this study was directly proportional to the difference of the square roots of the hydrogen partial pressures between the reaction and the permeate side. This is an ideal behavior for thick metallic membranes, which means the hydrogen permeation is controlled by the hydrogen diffusion in the membrane.10 For tubular membranes, the permeability of such a membrane [mol/ m‚s‚Pa1/2] can be derived from hydrogen permeation rate Q [mol/s] and hydrogen partial pressures of both sides Pr and Pp [Pa] by the following equation:

Q)

2πl0P h

(xPr - xPp)

ln(ro/ri)

(1)

where ri and ro [m] are the inner and the outer radii of the membrane tube and l0 [m] is the entire length. The permeability determined by using this equation is shown in Figure 2. The Arrehnius plot is found to be linear above 200 °C. Because the slope in this region gives an activation energy Ep of 6.4 kJ/mol, the permeability as a function of temperature is obtained as follows:

P h ) 3.0 × 10-8 exp(-Ep/RT)

(2)

In the low-temperature region, a broad peak in permeability due to the R-β phase transition11 is also found. 3.2. Decomposition of Methanol. 3.2.1. Measured Decomposition Rate. Methanol decomposition in the CCR mode was carried out mainly at pressures of 0.1 and 0.2 MPa and initial methanol concentrations of 20, 40, 60, 80, and 100%. Typical results are shown in Figure 3, where the molar ratio of the produced carbon monoxide to the feed gas, y, is plotted against L/N0, in which L [m] is the entire length of the catalyst layer and N0 [mol/m2‚s] is the feed flux of the gas mixture. This figure shows that y approaches the value of its initial concentration with increasing L/N0, indicating that the methanol conversion can reach completion at 220 °C if the reaction time is long enough. Because one of the aims of this study is to obtain higher conversions

Figure 2. Arrhenius plot of the hydrogen permeability of Pd91Ru6In3 alloy.

at lower temperatures by introducing the PMR, the decomposition was performed mainly at a temperature of 220 °C, where some improvement in the decomposition rate might be expected. From the slope of the plots in Figure 3, dy/d(L/N0), the reaction rate at the outlet of the reactor, r [mol/m3‚s], can be obtained. The initial rates r0, i.e., r at L/N0 ) 0, are well-consistent with each other without respect to the initial concentration. This feature gives us important information about the reaction mechanism. 3.2.2. Consideration on the Decomposition Kinetics. In terms of methanol decomposition, several rate equations have been proposed.12-15 Darby and Kamball14 empirically obtained the following equation for the reaction over a Fischer-Tropsch cobalt catalyst:

-

kPM dPM ) dt 1 + bPC

(3)

where PM and PC are partial pressures of methanol and carbon monoxide, respectively. On the other hand, Yasumori et al.15 carried out the reaction over a coiled nickel wire catalyst and found that the initial rate was expressed by

r0 )

aPM (1 + bxPM)2

(4)

All of the formulas proposed so far, which usually assume that the surface reaction on the catalyst is ratecontrolling, give initial rates strongly dependent on the methanol partial pressure. This is contradictory to the present results. The most possible formula based on surface reaction controlling is

r)

k(PM - PCPH2/K) 1 + KMPM

(5)

where k and KM, are a rate constant and an equilibrium adsorption coefficient and the subscripts M, H, and C indicate methanol, hydrogen, and carbon monoxide, respectively. The equilibrium constant for the methanol decomposition, K, calculated from the thermodynamic data is 1.06 × 1012 Pa2 at 220 °C. This rate equation provides an initial rate independent of PM if KM is large

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enough, which is consistent with the experimental results. The validity of this equation will be discussed later. One of the possible rate-controlling steps other than surface reaction is methanol adsorption onto the catalyst. In this case, assuming that the adsorption of any molecules except the reactant is negligible, the initial reaction rate for decompositions over catalysts should be directly proportional to the reactant partial pressure.16,17 However, this prediction is not in agreement with the present results. Therefore, it can be said that the overall reaction in our experiment is not controlled by the adsorption step either. The remaining possibility is that the desorption process of one of the products is rate-controlling. Assuming that the hydrogen desorption controls the overall rate and the adsorption of any other molecules is negligible, the reaction rate can be expressed as follows:

r)

k(xKPM/PC - PH)

Figure 3. Plots of y against L/N0 at various initial methanol concentrations.

(6)

1 + KHxKPM/PC

In the same way, if the desorption of carbon monoxide is rate-controlling, the reaction rate is given in the form

r)

k(KPM/PH2 - PC)

(7)

1 + KCKPM/PH2

Let us note that these two equations give initial rates independent of the methanol partial pressure as the hydrogen and the carbon monoxide partial pressure approach zero,15 which is consistent with the present results. 3.2.3. Determination of Rate Expression. To determine which process is rate-controlling, surface reaction, hydrogen desorption, or carbon monoxide desorption, eqs 5-7 are rewritten as follows:

Surface reaction: PM - PCPH2/K 1 KM ) + P r k k M

(8)

Hydrogen desorption:

x

xKPM/PC - PH ) 1 + KH r

k

Carbon monoxide desorption:

k

KPM PC

KPM/PH2 - PC 1 KC KPM ) + r k k P 2

(9)

(10)

H

According to these three equations, the experimental results carried out at atmospheric reaction pressure using 100% methanol are replotted in Figure 4. In contrast to parts a and b of Figure 4, the linearity between (KPM/PH2 - PC)/r and KPM/PH2 in part c is clearly seen to be well. It can, therefore, be concluded that the reaction in this study is controlled by the desorption of carbon monoxide produced. The constants k and KC are evaluated by linear least-squares fits to the data in Figure 4c, being 7.6 × 10-9 mol/m3‚s‚Pa and 5.7 × 10-8 1/Pa, respectively.

Figure 4. Correlations based on eqs 8-10 to determine the ratecontrolling step.

Using the obtained reaction rate equation, the methanol conversion can be reproduced by solving a differential equation expressing the mass balance in the reactor. The calculated results, shown in Figure 3 by

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Figure 5. Ratio of hydrogen to carbon monoxide in the effluent gas for the three different modes at reaction pressures of (a) 0.1 and (b) 0.2 MPa.

solid lines, agree with the experimental results, indicating that the rate equation was successfully determined. Moreover, the influence of two or more adsorption terms in the denominators of the rate equations 5-7 is also examined. In general, the denominator includes adsorption terms for all of the species in the mixture. For example, the denominator of eq 7 should be 1 + KCKPM/PH2 + KMPM + KHPH + KArPAr. However, the nonlinear least-squares fits of such an equation to the experimental data showed that, in this experimental range, the reaction rate was less sensitive to the last three terms than to the second term, leading to much error in KM, KH, and KAr determined by the fits. Taking account of the fact that eq 7 can reproduce the experimental data well enough, the last three adsorption terms do not seem to be essential. 3.3. Hydrogen Removal Effect on Methanol Conversion. The ratios of H2 to CO in the effluent gas in the three different modes are shown in Figure 5, where GHSV [1/h] is defined as a ratio of the gaseous methanol feed rate [m3(STP)/h] to the volume of the catalyst layer [m3]. Because one methanol molecule decomposes into one CO and two H2 molecules, the deviation from 2.0 means that H2 produced by the decomposition is removed from the reaction region through the membrane tube. In the CCR mode, the ratios are around 2.0. In the case of the two PMR modes, the ratios are less than 2.0 and decrease with decreasing GHSV, that is, increasing residence time of the reactant. It is obvious that the increase of the reaction pressure or the evacuation of the permeate side causes a large difference of the hydrogen partial pressures between both sides, namely, a large driving force for hydrogen permeation. Especially, in the PMR vacuum mode (0.1 MPa, GHSV ) 1.2 h-1) approximately 99% of hydrogen produced is removed from the reaction side. The improvement of methanol decomposition by hydrogen removal is shown in Figure 6. The most signifi-

Figure 6. Methanol conversion for the three different modes at reaction pressures of (a) 0.1 and (b) 0.2 MPa.

Figure 7. Conversion change at GHSV ) 10 h-1 with reaction pressure.

cant effect at a reaction pressure of 0.2 MPa can be seen at GHSV ) 4.2 h-1, where 24% larger conversion in the PMR vacuum mode than that in the CCR mode is achieved. It is also found that the higher the reaction pressure is, the greater the improvement effect becomes. The large conversion improvement at high reaction pressures (Figure 7) is due to not only an increase in the hydrogen permeation rate driven by the large pressure difference across the membrane, as shown in Figure 5, but also a decrease in the equilibrium conversion for the CCR. However, the improvement obtained is not large enough compared with the results reported elsewhere.4,8,9 Therefore, a mathematical model was solved by using the obtained permeability, eq 2, and the reaction rate, eq 7, where plug flow and steady state were assumed. As shown in Figure 8, the calculated conversion for the PMR mode is also larger than the experimental results. At such a low temperature as in this study, it is known that a decrease in permeability resulting from an inhibition of produced carbon monoxide often becomes critical:18 this might be responsible for the unsatisfied performance of the PMR. A more

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4. Conclusions

Figure 8. Conversion at a reaction pressure of 0.1 MPa calculated from the permeation rate and the reaction rate equation.

72q;1Methanol decomposition to carbon monoxide and hydrogen has been carried out by using a PMR, where a Pd91Ru6In3 alloy membrane tube of 0.2 mm thickness and 208 mm length was placed within a Pd/SiO2 catalyst layer. First, the kinetic analysis showed that the overall reaction was controlled by the desorption process of carbon monoxide from the catalyst surface. The promotion of the decomposition reaction by using the PMR was demonstrated at 220 °C using 100% methanol as a feed. The difference in the performance between the CCR and the PMR vacuum modes was found to increase with increasing reaction pressure. It was also found that hydrogen removal from the reaction region affected the amounts of byproducts. Literature Cited

Figure 9. Effect of the operation mode on selectivities to byproducts at GHSV ) 10 h-1.

detailed analysis including the inhibition will be described elsewhere. 3.4. Changes in Selectivity to Byproducts by Hydrogen Removal. Over all of the experimental conditions, the selectivity to CO was larger than 94%. The byproducts were CO2 and a small amount of CH3OCH3, CH4, HCOOCH3, and H2O. HCHO was detected only at a reaction pressure of 0.1 MPa and not at elevated pressures. The total selectivity to byproducts at high pressures, as shown in Figure 9, tended to be reduced by the hydrogen removal. Some of these byproducts might be considered to be produced by the following side reactions:

2CH3OH S CH3OCH3 + H2O

(11)

CO + H2O S CO2 + H2

(12)

CH3OH S HCHO + H2

(13)

2HCHO S HCOOCH3

(14)

During the hydrogen removal, because the main reaction of the methanol decomposition is promoted and the methanol partial pressure decreases, the forward reaction in eq 11 is depressed, thereby resulting in a decrease in CH3OCH3 selectivity as shown in Figure 9. This also leads to a decrease in H2O, which might restrict the CO2 production in eq 12. On the other hand, the selectivity to HCOOCH3 in Figure 9 increases by the hydrogen removal. This is probably caused by an increase in the amount of HCHO produced by the hydrogen removal according to eq 13, followed by rapid conversion to HCOOCH3 due to the instability of HCHO.

(1) Uemiya, S.; Matsuda, T.; Kikuchi, E. Aromatization of Propane Assisted by Palladium Membrane Reactor. Chem. Lett. 1990, 1335. (2) Champagnie, A. M.; Tsotsis, T. T.; Minet, R. G. A High Temperature Catalytic Membrane Reactor for Ethylene Dehydrogenation. Chem. Eng. Sci. 1990, 45, 2423. (3) Zaika, Z. D.; Minet, R. G.; Tsotsis, T. T. A High Temperature Catalytic Membrane Reactor for Propane Dehydrogenation. J. Membr. Sci. 1993, 77, 221. (4) Edlund, D. J.; Pledger, W. A. Thermolysis of Hydrogen Sulfide in a Metal-Membrane Reactor. J. Membr. Sci. 1993, 77, 255. (5) Cacciloa, G.; Giordano, N.; Restuccia, G. Cyclohexane as a Liquid Phase Carrier in Hydrogen Storage and Transport. Int. J. Hydrogen Energy 1984, 9, 411. (6) Gretz, J.; Baselt, J. P.; Ullmann, O.; Wendt, H. The 100 MW Euro-Quebec Hydro-Hydrogen Pilot Project. Int. J. Hydrogen Energy 1990, 15, 419. (7) Wood, B. J. Dehydrogenation of Cyclohexane on a HydrogenPorous Membrane. J. Catal. 1968, 11, 30. (8) Itoh, N. A Membrane Reactor Using Palladium. AIChE J. 1987, 33, 1576. (9) Ali, J. K.; Baiker, A. Dehydrogenation of Methylcyclohexane to Toluene in a Pilot-Scale Membrane Reactor. Appl. Catal. A 1997, 155, 41. (10) Shu, J.; Grandjean, B. P. A.; van Neste, V.; Kaliaguine, S. Catalytic Palladium-Based Membrane Reactors: A Review. Can. J. Chem. Eng. 1991, 69, 1036. (11) Aoki, K.; Yokoyama, S.; Kusakabe, K.; Morooka, S. Preparation of Supported Palladium Membrane and Separation of Hydrogen. Korean J. Chem. Eng. 1996, 13, 530. (12) Quinn, D. F.; Taylor, D. The Decomposition of Formic Acid and Methanol on Copper-Nickel Alloys. J. Chem. Soc. 1965, 5248. (13) Steinbach, F.; Spengler, H. J. Temperature Jump Method in Molecular Beam Relaxation Spectrometry Applied to Catalytic Decomposition of CH3OH on Polycrystalline Ni Foil. Surf. Sci. 1981, 104, 318. (14) Darby, P. W.; Kamball, C. Investigation of Reactions Along the Catalyst Bed in Flow Systems by Vapour-Phase Chromatography. Trans. Faraday Soc. 1957, 53, 832. (15) Yasumori, I.; Nakamura, T.; Miyazaki, E. Catalytic Decomposition of Methanol over Nickel Wire. Bull. Chem. Soc. Jpn. 1967, 40, 1372. (16) Hougen, O. A.; Watson, K. M. Solid Catalysis and Reaction Rates. Ind. Eng. Chem. 1943, 35, 529. (17) Itoh, N.; Tanabe, H.; Shindo, Y.; Hakuta, T. Kinetic Analysis of Cyclohexane Dehydrogenation over Platinum Catalyst. Sekiyu Gakkaishi 1985, 28, 323. (18) Amano, M.; Nishimura, C.; Komaki, M. Effects of High Concentration CO and CO2 on Hydrogen Permeation through the Palladium Membrane. Mater. Trans., JIM 1990, 31, 404.

Received for review July 13, 1998 Revised manuscript received October 20, 1998 Accepted November 4, 1998 IE980456D