225
Ind. Eng. Chem. Prod. Res. Dev. 1983, 22, 225-228 Emerson, W. S. I n "Organic Reactions", Adams Jr., R., Ed.; Wliey: New York. 1948; Voi. 4, 6 1 7 4 . Franckaerts, J.; Froment, 0. F. Chem. Eng. Scl. 1984, 79, 807. Himmelblau, D. M. "Process Analysis by Statlstlcal Methods"; Wiley: New York, 1970; p 194. Issolre, J.; Van Long, C. Bull. SOC.Chlm. F r . 1960, 2004. Klaus, R.; Rippin, D. W. T. 12th Symposium on Comp. Appi. in Chemical Engineerlng, Montreux, 1979; p 155. Kliger, G. A.; Larutina, L. F.; Fridman. R. A.; Kryukow, Y. 6.; Bashkirov, A. An.: Snamvskii. Y. S.: Smirnova. R. M. Klnet. Kate/. 1975a. 76. 660. Kliger,'G. A.,Lazutina, L. F.; FrMman, R. A.; Kryukov, Y. 6.; Bashkirov, A. N.; Snagovskii, Y. S. Klnet. Katal. 1975b. 76, 665. Kliger, G. A.; Glebov, L. S.; FrMman, L. A.; Vytnova, L. A.; Bashkirov, A. N. Klnet Katal. 1978, 79, 615. Madon, R. J.; O'Conneii, J. P.; Boudart, M. AIChE J. 1978, 2 4 , 904. March, J. "Advanced Organic Chemistry: Reactions, Mechanisms, and Structure"; McGraw-Hill: New York, 1968. Miyamoto, A.; Ogino, Y. J . Catal. 1975, 3 7 , 133. Murchashi, S. I.; Hirano, T.; Yano, T. J . Am. Chem. SOC. 1978, 700, 348. Nystrom, R. F.; Brown, W. G. J . Am. Chem. SOC. 1947, 69, 1197.
Okamoto, Y.; Imanaka, T.; Teranishi, S. Bull. Chem. SOC.Jpn. 1972, 4 5 , 3207. Ono, Y.; Ishida, H. J . Catel. 1981, 7 2 , 121. "Organlkum"; VEB Deutscher Verlag der Wissenschaften: Berlin, 1977. Peiosa, A.; Moresl, M.; Mustachi, C.; Soracco, B. Can. J . Chem. Eng. 1979, 5 7 , 159. Rau, R. C. I n "Encyclopedia of X-Rays and y-Rays"; Clark, G. L., Ed.; ReinhoM: New York, 1963; p 184. Ryknder, P. N. "Catalytic Hydrogenation over Platinum Metals"; Academic Press: New York, 1967. Schwoegler. E. J.; Adkins, H. J . Am. Chem. Soc.1939, 67, 3499. Segal, E.; Madon, R. J.; Boudart, M. J . Cafal. 1978, 5 2 , 45. Stull, D. R.; Westrum, E. F.; Sinke, G. C. "The Chemical Thermodynamlcs of Organic Compounds"; Wiiey: New York, 1969. Vorob'ev, A. M.; Evseeva, G. V.; Zenkevich, L. V. Zh. Flz. Khlm. 1973, 47, 2708.
Received for review July 30, 1982 Accepted November 18,1982
Kinetics of the Hydrogenation of Carbon Dioxide over Supported Nickel Jeng H. Chlangt and Jack R. Hopper" Chemical EngIneerlng Department, Lamar Unlverslfy, Beaumont, Texas 777 70
The kinetics of the hydrogenation of carbon dioxide to methane was studied using a 58% nickel catalyst supported on kieselguhr. The volume percentages of carbon dioxide and hydrogen in feed mixtures were varied from 20 to 30% and 67 to 80%, respectively. A total pressure range of 100 to 250 psig and a temperature range of 530 to 605 O F was covered. Experimental data were correlated with the power rate model to give the following relationship: rCH,= 1.19 X 10' e ~ p [ - 1 4 6 0 0 / R T ] P , ~ ~ ~ ~ P ~ ~ ~ ' ' .
Introduction Recent interest in the methanation reaction is a result of this reaction being required as the final step in the production of substitute natural gas (SNG) from coal. The practical significance of this process will no doubt hinge on the creditability predicted for the severe shortages of natural gas for the imminent future (Hottel and Howard, 1971). The primary reaction for the methanator is the hydrogenation of carbon monoxide to form methane and water, and considerable effort has been expended in the study of this primary reaction using various catalysts. Most of the previous work has recently been reviewed (Vannice, 1976; Mills and Steffgen, 1973; Vlasenko and Yuzefovich, 1969). Although the primary reaction is the methanation of carbon monoxide, the shift reaction CO + H2O + COZ + H2 (1) and the methanation of carbon dioxide C 0 2 + 4H2 CHI 2H20 (2) are two secondary reactions which could be highly significant. A n analysis of the methanation reactions will no doubt be incomplete without an understanding of the contribution of these reactions. In view of the potential
+
'Alpha Chemical Laboratories, 299 Lesmill Rd., Toronto, Ont. M8VlB6. Canada. 0196-4321/83/1222-0225$01.50/0
significance of the methanation of carbon dioxide on the overall methanation reaction, this study was undertaken.
Previous Work The effect of carbon dioxide on carbon monoxide methanation which results indirectly through the shift reaction and directly fiom the hydrogenation of carbon dioxide has been recognized since the early investigations of the hydrogenation of carbon monoxide. Akers and White (1948) realized the importance of C 0 2 hydrogenation in one of the earliest kinetic studies of carbon monoxide, but they did not have sufficient data to make an accurate analysis of C 0 2 rates. Binder and White (1950) acknowledged the significance of C 0 2 by studying the rate of methanation of C 0 2alone and observed the methanation rate for C 0 2to be two orders of magnitude less than that for CO. This work at the University of Michigan was continued by Dew et al. (1955). Solc (1962) and Pour (1969) studied the kinetics of carbon dioxide hydrogenation on a chromium-nickel catalyst using a large excess of hydrogen in the flow system feed gas. Both reported the rate to be half-order with respect to carbon dioxide. Vlasenko et al. (1961,1965) have reported on hydrogenation studies using pure CO and C02 and mixtures of the two. However, carbon oxide concentrations were less than 3% for the purpose of studying the purification of hythe rate drogen streams. In a study using pure C02 (lsl), of methanation was observed to be first order with respect to COPand no appreciable adsorption of C02was observed. Methane and water did not influence the rate. Vlasenko 0 1983 American Chemical Society
226
Ind. Eng. Chem. Prod. Res. Dev., Vol. 22, No. 2, 1983 T E M P E R ATUR E
I
A I R OUT
\
R
1 Figure 1. A simplified flow diagram of the reaction system.
et al. (1961,1965) also observed that carbon dioxide does not affect the rate of CO methanation but carbon monoxide greatly reduces the rate of carbon dioxide methanation. Van Herwijnen et al. (1973) examined the kinetics of the methanation of both carbon monoxide and carbon dioxide on nickel catalysts. Similar to Vlasenko’s work, it was done at low concentrations of the oxides, and similar observations for CO, C02, and products were made. Saletore and Thomson (1977) reported on a study including all five components present in the methanation reactions. The presence or absence of C02 and water in the system feed was demonstrated to have a significant role in methanation kinetics and it was proposed that C o t methanation might actually proceed via the water-gas shift reaction. At the higher concentrations of carbon oxides, water reduced the rate of both the shift reaction and CO methanation, but methane had little effect. An increase in carbon dioxide increased the CO methanation rate slightly but decreased the shift reaction rate significantly. Thomson and Murphy (1979) more recently extended this initial study and observed that C02 is directly hydrogenated to methane rather than indirectly through the shift reaction as originally speculated. Moore (1977) made an extensive study of the hydrogenation of carbon monoxide with some experiments including the addition of carbon dioxide, methane, and water to the feed. He observed that an increase in carbon dioxide concentration increased the rate of disappearance of carbon dioxide and simultaneously a slight increase in the rate of formation of methane; this was speculated to be due to the formation of methane from carbon dioxide at a rate faster than the formation of C 0 2 in the shift reaction. Methane concentration up to 57% had little influence on the CO methanation rate, while the addition of water caused a significant decrease in methane formation but only a slight increase in C02 formation. Experimental Methods Apparatus. The kinetic studies were conducted in a continuous-flow tubular reactor system previously used by Norman et al. (1976). A simplified flow diagram of the reaction system is given in Figure 1. The system consisted of a gas feed system, a fluidized sand bath to maintain constant reactor temperature, a back-pressure regulator in the effluent line, and a wet-test meter for product gas measurement. Hydrogen and carbon dioxide were regulated independently before mixing prior to entering the reactor. The reactor was constructed of 1/4-in.316 stainless steel tubing, and normally about 0.15 g of catalyst was contained in the reactor between two 15-km in-line filters.
POROUS FRIT
1~
Figure 2. A cross-sectional view of the reactor and the heating assembly. Table I. The Surface Properties of Ni-104 Catalyst metal content, wt % support original pellet dimension, in. total surface, m’/g pore volume, cm3/g mean pore (radius), .A metal surface area, m’/g of cat. metal surface area, m’1g of metal average crystallite size, A dispersion, o/o
58% Ni kieselguhr ‘1s
x
‘/s
144.6 0.114 13.1 39.2 61.6 83.2 11.65
A cross-sectional view of the reactor in the fluidized sand bath is shown in Figure 2. Product analysis was performed with a modified Micro-Tek temperature-programmed, dual-column gas chromatograph. Peak area integration was performed by an Infotronic Digital Integrator Model CRS-208 with output signal to a printer. Helium was used as the carrier gas (20 cm3/min) for a 10-ft GC column packed with 50-80 mesh Porapak Type Q and a 15-ft column packed with 10% 13x pellets and 90% 5A pellets (80-100 mesh) Linde Molecular Sieves. The column and detector were held a t 140 “C. Complete descriptions of the equipment and procedure have been presented earlier (Chiang, 1978). Materials. Harshaw Reduced Nickel Catalyst No. 104, Lot No. 786 containing 58% nickel supported on kieselguhr with a granule size of 0.07-0.1 mm was used for all experiments. Surface properties for this catalyst were determined by Luna (1976) and are shown in Table I. Hydrogen gas of 99.5% purity and caybon dioxide gas of 99.86% purity were used as received. Experimental Results The rate of hydrogenation of carbon dioxide to methane was measured over a range of temperature from 530 to 605 O F and a range of total pressure from 155 to 250 psig. A range of space time from 0.03 to 0.2 (h) (g of cat.)/ft3 and a range of volume percent carbon dioxide in the feed from 17 to 33% were used. The catalyst was reduced with hydrogen a t 800 O F and 500 psig for 48 h after initially being put in the reactor. The catalyst was reduced at a temperature greater than any used in the kinetic experiments to minimize thermal effects of catalyst activity. A steady-state condition in catalyst activity was reached before each test by operating at least 24 h with a 25% carbon dioxide in hydrogen mixture at a pressure of 300 psig, a temperature of 500 OF, and a space velocity of 17.1 (ft3)/(h)(gof cat.). To provide consistent kinetic data a reference set of reaction conditions was used to standardize the catalyst activity and
Ind. Eng.
Chem. Prod. Res. Dev., Vol. 22, No. 2, 1983 227
Table 11. Activation Energies for the Hydrogenation of Carbon Dioxide
59% Ni-88 59% Ni-88
541-747 500-752 356-572
reduced NiO o n kieselguhr reduced NiO-Cr,O, Ni-Cr,O,
2.1-30
1
1
I
I
SPACE TIME
6-89 25 or excess not available 1 less than 1
320-356 320-428 392-414 257-617 530-605
33.6% G-65 51.9% Ni-Cr,O, 58% Ni-104 I
reactor press., a t m CO, concn, %
reaction temp, "F
cat.
1 6-18 I
1.4 0.17-0.51 20-33
act. energy, kcal/mol
ref
13.1-13.9 16 o r 30 13.1
Dew e t al. (1955) Binder and White (1950) Nicolai e t al. (1946)
20.6 13.8,17.3 25.3 13.9 14.6
Solc (1962) Pour (1969);Muller e t al. (1968) Van Herwijnen e t al. (1973) Vlasenko et al. (1961,1965, 1969) this work
I
hr-gm/f t
Figure 3. Conversion of COz as a function of space time at variable total pressure, Hz/Co2= 41: (0) 250 psig; (0) 200 psig; (A)155 psig; ( 0 )100 psig.
SPACE T I M E
hr-gm/fts
Figure 4. Conversion of C02a~ a function of space time at variable hydrogen to carbon dioxide mole ratio at 250 psig, 550 O F ; (0) Hz/ COP = 5:l;(0) Hz/COz = 4:l;(A)HZ/COz = 3:1; (0)Hz/C02 = 2:l.
correct for any deviations in catalyst activity from this standard. External transport resistances were minimized by using superficial mass velocities in the range of 183-1179 (lb,)/(h)(ft2) which were calculated to be wellabove the necessary 40 lb,/(h)(ft2) at 700 OF. To minimize intraparticle transport effects, particle sizes of 140-200 mesh were used. This should be sufficient to eliminate these effects since it was previously shown for the carbon monoxide hydrogenationthat particle sizes up to 40 mesh could be used at 700 O F (Moore, 1977). The rate of appearance of methane was calculated using the data for the conversion of carbon dioxide as a function of reciprocal space velocity as follows
x = f(W/F) ~
C
=H dX/d(W/I;? ~
(3)
= kPH:PCO;Y
v
u 1.70
1.74
1.78
1.82
(4)
A typical plot of conversion as a function of reciprocal space velocity is shown in Figure 3 for total pressures of 100,155,200,and 250 psig. A similar plot for H2/C02rates of 2:1, 3:1, 4:1, and 5:l is shown in Figure 4. A regression analysis was used to obtain the functional relationship from which the rate was calculated. Most of the rate data were obtained for carbon dioxide conversion of 10 to 20% and all of the conversion values were less than 45%. Carbon monoxide concentration was followed closely to detect the possibility of formation from the reverse shift reaction. A large number of the experiments had no measurable carbon monoxide in the product. Only at H2/C02 ratios of 2:l was carbon monoxide measurable and it was less than 2.0% in all cases. Discussion A Simple Power Law Correlates the Data. A power rate law model of the form rCHd
0.5
(5)
(where rcK = rate of formation of methane (g-mol)/(day)(g
I/T, 103/.K Figure 5. Arrhenius plot for the power-law rate model.
of cat.), k = specific reaction rate constant (g-mol)/(day)(g of cat.)(atmIx+Y,Pco2,PHz = partial pressure of hydrogen and carbon dioxide, atm, and x , y = order of reaction parameters for hydrogen and carbon dioxide) has been used to initially correlate the rate data. The effect of temperature on the specific reaction rate constant, k , was correlated with the Arrehnius equation
k = A exp(-E/RT)
(6)
where E = molal energy of activation (cal/g-mol), A = frequency factor (g-mol)/(day)(g of cat.)(atm)x+Y,R =
universal gas constant cal/(g-mol K), and T = absolute temperature, K. Values of x and y were found to be 0.21 and 0.66 for the order of reaction with respect to hydrogen and carbon dioxide, respectively. Values for the rate constant at the various temperatures are shown in the Arrhenius Plot of Figure 5. An activation energy of 14.6 kcal/mol was obtained from this relationship
Ind. Eng. Chem. Prod. Res. Dev., Vol. 22, No. 2, 1983
228
D
2 i_"s z
lr
i
41 t
01 0 0
002
004
006
,
010
008
S P A C E T I ME
Oil
0 12
1
014
t-
\ 4-1 ! 6
hr-gm/ft3
Figure 6. Comparison of experimental and calculated rate data at H2/CO2= 4:1, 250 psig, 570 OF: (-1 exptl; (- - -) calcd.
i 0 . 5L
Table 111. Adsorption Model Parameters rate constant, k
frequency factor 1 4 0 g-mol/(day)term ( g of cat.)(atm)' activation energy, 4451 E,, kcal/gmol heat of adsorption, A H , kcalig-mol
L
L
1.70
ads constants model constants
KH,
KCO,
3984
and the value for the frequency factor was calculated to be 1.19 X lo6 (g-mol)/(day)(g of ~at.)(atm)O.~'.A comparison of this value for the energy of activation with values previously reported is given in Table I1 and is seen to be in good agreement with these data which cover a wide range of conditions. Thus the simple power law rate model can be represented as rCH4= 1.19 X lo6 exp[-14600/RT]P~20.~*Pc020~~~ (7) This equation fits all of the data with an average deviation of 6.93%. A graphical illustration of the agreement between this model and one set of data is shown in Figure 6. Two-step irreversible uniform and nonuniform surface reaction models (Boudart, 1972; Vannice, 1975) were derived and evaluated, but none of these models fits the experimental data with a physical reality. A Hougen-Watson-Langmuir-Hinshelwood adsorption type model of the form kPH~PCO+
where k = ko exp(-E,/RT)
Kco2 = Kcozo exp(-AHco,/RT)
L 1.74
178
l
h
4
1.82
I / T , i0Y"K
7.2 x 15' 0.018 L/atm L/atm 17823
9
4-2 0
(11)
was also used to correlate the rate data. The value for the rate constant k and the adsorption constants KH2and KCO, were determined separately at each temperature. The
Figure 7. Arrhenius plot for the adsorption model: (0) k ; (0) KH2; (v)Kco2.
average deviation for these data is 4.61%. The results for this model are presented in the Arrhenius plot in Figure 7. The values for the parameters in this model are given in Table 111. Although the parameters of this model have physically realistic values the model is considered at best to be a mathematical representation of the data and as such the power law model is more valuable because of the simplicity.
Acknowledgment The financial support from the National Science Foundation for Grant No. ENG 74-04191 A01 is gratefully acknowledged.
Literature Cited Akers, W. W.; White, R. R. Chem. Eng. Rog. 1948, 4 4 , 553. Binder, G. C.; White, R. R. Chem. Eng. Prog. 1950, 4 6 , 563. Boudart, M. AIChE J. 1972. 18, 465. Chiang, J. H. M.E.S. Thesis, Lamar University, Beaumont, TX, 1978. Dew, J. N.; White, R. R.; Sllepcevich, C. M. Ind. Eng. Chem. 1955, 4 7 , 140. Hottel, H. C.; Howard, J. 8. "New Energy Technology-Some Facts and Assessments"; The MIT Press: Cambridge, MA, 1971. Luna, V. M.E.S. Thesis, Lamar University, Beaumont, TX, 1976. Mills, G. A.; Steffgen Catal. Rev. 1973, 8, 159. Moore, G. E. M.E.S. Thesls, Lamar University, Beaumont, TX, 1977. Muller, J.; Pour, V.; Regner, A. J. Catal. 1968, 7 1 , 326. Nicolal, J.; D'Hont, M.; Jungers, J. C. Bull. SOC.Chlm. Belg. 1948, 55, 160. Norman, 0.H.; Shlgemura, D. S.;Hopper, J. R. I n d . Eng. Chem. Prod. Res. Dev. 1976, 15, 41. Pour, V. Collect. Czech. Chem. Commun. 1969, 3 4 , 45. Saktore, D. A.; Thomson, W. J. Ind. Eng. Chem. Process Des. Dev. 1977, 16, 70. Solc, E. Collect. Czech. Chem. Common. 1962, 27. 2621. Thomson, W. J.; Murphy, M. L., paper presented at 86th National AIChE Meeting, Houston, TX, Apr 2, 1979. Van Herwijnen, T.; Van Doesburg, H.; Dejong, A. J. Catai. 1973, 28, 391. Vannice, M. A. J. Catal. 1975, 37, 462. Vannice, M. A. Catal. Rev.-Sci. Eng. 1976, 14, 153. Vlasenko, V. M.; Russov, M. T.; Yuzefovich, G. E. Kinef. Catal. 1961, 2(4), 476. Vlasenko, V. M.; Russov, M. T.; Yuzefovich. G. E. Kinet. Catal. 1965, 6(5), 849. Valsenko, V. M.; Yuzefovich, G. E. Russ. Chem. Rev. 1989, 3 8 , 728.
Receiued for reuiew July 30, 1979 Revised manuscript received December 7 , 1982 Accepted February 14, 1983
