Ind. Eng. Chem. Res. 1987,26, 2461-2466 = air kinematic viscosity = standard deviation 7 = particle mean residence time in the reactor humidity (mass fraction) at the reactor inlet xi, X e = V,
u
and outlet, respectively
Literature Cited Buchmann, R. M&. Dissertation, Ben Gurion University of the Negev, Beer Sheva, Israel, 1985. Cheri, S. S. AIChE J . Symp. Ser. 1971,116,67. Elperin, I. T. Znz. Fiz. Zhur 1961,6,62. Elperin, I. T.; Tamir, A. Israeli Patent 66 162,1982.
2461
Kitron, A. M.Sc. Dissertation, Ben Gurion University of the Negev, Beer Sheva, Israel, 1984. Luzzatto, K.; Tamir, A.; Elperin, I. AIChEJ. 1984,30(40), 600. Shalmon, B. BSC.Dissertation, Ben Gurion University of the Negev, Beer Sheva, Israel, 1985. Tamir, A.; Elperin, I.; Luzzatto, K. Chem. Eng. Sci. 1984,39(1),139. Tamir, A.;Luzzatto, K. AZChEJ. 1985,31(5), 781. Tamir, A,; Luzzatto, K.; Sartana, D.; Surin, S. AIChEJ. 1985,31(10), 1744. Tamir, A.; Kitron, A. Chem. Eng. Commun. 1987,50, 241.
Receiued for review September 9, 1986 Revised manuscript received July 23, 1987 Accepted August 17,1987
Solubility of COz in Aqueous Methyldiethanolamine and N,N-Bis(hydroxyethy1)piperazine Solutions Amitabha Chakma and Axel Meisen* Department of Chemical Engineering, The University of British Columbia, Vancouver, British Columbia V6T 1 W5, Canada
The solubility of carbon dioxide in 4.28 and 1.69 M aqueous methyldiethanolamine (MDEA) and 0.287 and 0.115 M NJV-bis(hydroxyethy1)piperazine (BHEP) solutions has been determined experimentally. The experimental conditions were 100-200 OC for MEDA, 40-180 “Cfor BHEP, and C02partial pressures of 172-4929 W a . Mathematical models for predicting COz solubility in MDEA and BHEP solutions are presented. The model correlations agree well with present and previously published data. Aqueous methyldiethanolamine (MDEA) solutions are finding good acceptance in industry for the selective removal of H$ from light hydrocarbon gases also containing COz (Frazier and Kohl, 1950; Vidaurri and Khare, 1977; Kohl and Riesenfeld, 1985). A further advantage of MDEA over other amines is that it does not degrade readily; i.e., it does not enter reactions which are irreversible under normal process conditions (Blanc et al., 1981). However, Chakma and Meisen (1985) recently reported that MDEA reacts fairly rapidly with carbon dioxide at elevated temperatures and C02 partial pressures forming, amongst other compounds, NJV-bis(2-hydroxyethy1)piperazineor “BHEP”. The latter compound is also produced during the degradation of diethanolamine (DEA), which is a widely used industrial solvent for acid gases (Kohl and Riesenfeld, 1985), and may occur in concentrations of several percent in industrial solutions (Smith and Younger, 1972; Kennard and Meisen, 1985). Since BHEP is not corrosive toward mild steel and, like amines, can absorb C02,it is often thought to be tolerable in industrial solutions. The solubility of carbon dioxide in BHEP is not presently known. Jou et al. (1982) have reported C02solubilities in MDEA solutions up to 120 “C. However, the MDEA degradation reactions become more important at even higher temperatures, and the degradation rates can only be predicted provided the COz solubilities are known. The primary objectives of this paper are therefore to provide data on the solubility of COz in MDEA solutions at temperatures up to 200 “C and in BHEP solutions at temperatures ranging from 40 to 180 “C. In addition, mathematical models are presented for predicting the solubilities. The models, which are extensions of the work by Kent and Eisenberg (1976) and Jou et al. (1982), are easier to use than theirs since they do not depend on graphical data and can be solved completely by computer. 0888-5885/87/2626-2461$01.50/0
Experimental Apparatus and Procedure The experimental apparatus and procedures are similar to those described by Kennard and Meisen (1985). Aqueous MDEA and BHEP solutions of the desired concentration were prepared by mixing distilled water with 99+% pure MDEA and BHEP (supplied by Aldrich Chemicals, Milwaukee, WI). The purity of MDEA and BHEP was confirmed by gas chromatography. A 500-mL, high-pressure, stainless steel bomb was filled with C02and weighed with an electronic balance (Type P2000, Mettler, Switzerland) to within f O . l g. A 600-mL stainless steel autoclave (Model 4560, Parr Instrument Co., Moline, IL) was charged with 200 mL of the prepared MDEA or BHEP solution and closed. The autoclave was then heated to the desired temperature. Its contents were constantly stirred and the temperature was kept at the desired value by means of a controller (Model 4831EB, Parr Instrument Co.). The pressure of the solution was read from the gauge attached to the autoclave. The C02-filledbomb was then connected to the autoclave. Equilibrium was usually reached within less than half an hour as noted from any changes in the pressure reading. However, the system was allowed to equilibrate for at least 4 h before the 500-mL bomb was disconnected from the autoclave. The equilibrium pressure was then recorded. The bomb was reweighed and the mass of C02introduced into the autoclave was calculated from the weight change. Theory C 0 2 Solubility. The solubility of C02 was calculated as follows. First, the amount of C02 in the vapor phase was determined by assuming that the phase obeyed Dalton’s law. The partial pressure of C02 in the vapor phase was therefore calculated from the difference between the final equilibrium pressure and the solution pressure before the COPwas introduced. The number of moles of C02 in 0 1987 American Chemical Society
2462 Ind. Eng. Chem. Res., Vol. 26, No. 12, 1987
the vapor phase was then found by using the B-W-R equation of state (Benedict et al., 1951), the partial pressure of COP,the temperature and the volume of the vapor phase which was calculated from the known internal volume of the autoclave, and the volume of the MDEA or BHEP solution. The latter was determined at room temperature and its thermal expansion was neglected. Furthermore, the effect of vaporization on liquid volume and liquid concentration was found to be negligible. The BW-R constants for COz were obtained from Reid et al. (1977). The mass of CO, dissolved in the MDEA or BHEP solution was then found by subtracting the mass of CO, in the vapor phase from the mass of COz introduced into the autoclave. Mathematical Modeling of COz Solubility in MDEA. In the absence of degradation reactions, the equilibrium in the MDEA-CO2-H20 system is governed by the following ionic reactions: RzNCH3+ COz + HzO + R2NCH3H++ HC03- (1) H 2 0 + COP+ H+ + HC03(2) H,O ==H+ + OH(3) HC03- + H+ Henry's law may be written as
+ C032-
(4)
(5) where PCO,and HCo,denote the partial pressure of carbon dioxide in the gas phase and Henry's law constant, respectively. Molar concentrations in the liquid phase are indicated by square brackets, i.e., [ I. The corresponding equilibrium expressions are [RZNCH3H+][HCOS-] K,' = (6) [RzNCH,I [ c o d K3 = [H+1[HC03-1/ [Cozl (7) K4 = [H+][OH-] (8)
K5 = [H+][C032-]/[HC03-] (9) The equilibrium constant defined in eq 6 is denoted by K< to distinguish it from K1 used by Kent and Eisenberg (1976). In addition, the total molar and charge balances must be satisfied: [MDEA] = [RZNCH,] + [R2NCH3Hf] (10) [MDEAIa = [COZ] + [HCOS-] + [CO,2-] (11) [R,NCH3H+] + [H+] = [OH-]
+ [HC03-] + 2[CO?-]
(12) where [MDEA] and a denote the total MDEA concentration and CO, loading in the solution, respectively. Ekpmtions 6-12 can be written to express the COz partial pressure as a function of acid gas loading in the solution:
K1' [MDEA] K3 + K1' [H'] K5 + [H+l (14) Provided the CO2-MDEA-H20 system behaves ideally, the Henry's law and equilibrium constants defined by eq 5-9 should only be functions of temperature. Kent and
Eisenberg (19761, who investigated the CO2-MEA-H20 and COP-DEA-H,O systems, found this to be the case and presented empirical expressions for calculating the constants. However, for the C02-MDEA-H20 system, Jou et al. (1982) were only able to predict their experimental measurements satisfactorilyby assuming that Hcq, K3,K4, and K5 depended on temperature. (They used Kent and Eisenberg's (1976) empirical expressions to calculate these constants.) However, for the equilibrium constant governing the main amine reaction (i.e., for K l ) ,they had to postulate a dependence on temperature, acid gas loading ( a ) ,and total amine concentration. This modification is equivalent to including all system nonidealities in the expression for K,. Deshmukh and Mather (1982) have developed a more rigorous model, requiring calculation of activity and fugacity coefficients. While they determined the fugacity coefficients from the Peng-Robinson (1976) equation of state, the activity Coefficientswere calculated by using the extended Debye-Huckel expression given by Guggenheim (1935). The use of the latter expression requires a knowledge of interaction parameters for all pairs of species in solution. Since these parameters were neither available nor could be obtained within the scope of this work, the model of Deshmukh and Mather (1982) could not be used here. A method similar to that of Jou et al. (1982) was developed for this work. In particular, Kent and Eisenberg's (1976) expressions for Hco,, K3,K4, and K5 were chosen. Data obtained from the present study and Jou et al. (1982) were used to determine K< as a function of temperature, [CO,], and [MDEA]. Using [CO,] is preferable to a since the former can be calculated directly from Pco,and Hco,. Thus, the equilibrium constant expressions are K,' = exp(92.421453- 1.49081486 x 1 0 - T + 40.847708/T - 14.031652 In (2') 9.8778738 X 10-2[C02]+ 0.18275505 In [CO,] + 3.9862282[MDEA] - 12.715421 In [MDEA]) (15)
K3 = exp(-241.818 + 298.253 X 103T1- 148.528 x 106T2+ 332.648 X 108T3- 282.394 X 1010T4) (16) K4 = exp(39.5554- 987.9 X 102T1+ 568.828 X 105T2- 146.451 X 108T3+ 136.146 X l O l o T 4 ) (17) K , = exp(-294.74 + 364.385 X 103T1- 184.158 X 106T2+ 415.793 X 108T3- 354.291 X 1010T4)(18) HcOz= exp(22.2819 - 138.306 X 102T1+ 691.346 x 104T2- 155.895 X 107T3+ 120.037 X 109T41/7.50061 (19) Mathematical Modeling of COPSolubility in BHEP. The equilibria for the CO2-BHEP-H20 system are similar to those of the C02-MDEA-H20 system. Equation 1 needs to be replaced by [BHEPH+] + [H+] + [BHEP] (20) and eq 6 becomes Kbl = [H+][BHEP]/ [BHEPH'] (21) The mass and charge balance equations are B = [BHEP] + [BHEPH+] (22) Ba = [COz] + [HC03-] + [C03'-]
(23)
[BHEPH+] + [H+] = [OH-] + [HC03-] + 2[C032-] (24)
The following equations can then be derived to relate the
Ind. Eng. Chem. Res., Vol. 26, No. 12, 1987 2463
?H y E
i
-
2:
8-
10.
3
571ff
3
10.
CO2 PARTIAL PRESSURE ( k P a )
3
5 7
1v
3
5 7 104
COZ PARTIAL PRESSURE ( k P a )
57104
Figure 1. Comparison of the present C02solubility data in 4.28 M MDEA solutions with Jou et al.'s data and model predictions as a function of C 0 2 partial pressure (temperature, 100 "C).
j
Figure 2. Present experimental and predicted (solid lines) C02 solubilities in 4.28 M MDEA solutions as functions of C 0 2 partial pressure and temperature.
C02 partial pressure to acid gas loading at a given temperature and BHEP concentration:
[H+] =
+ [H+] K1 + [H+] + B Kbl
The following empirical correlation for Kbl was obtained by fitting the experimental data: + Kbl = exp(433.263099 - 2.8134775 X 35.949793/T 10.472318(1n T )+ 2.6761498[CO21 0.75472578(1n [CO,]) - 7.3049642Bl (27) Equations 13 and 14 together with eq 15-19 can be solved by using nonlinear equation solvers to calculate the solubility of carbon dioxide, a, provided the MDEA concentration, solution temperature, and CO, partial pressure are given. Unlike the procedure of Jou et al. (1982),our model allows the solubility calculations to be performed directly without an initial estimate of solubility for the calculation of K1.Furthermore, it is not restricted to any particular solution concentration within the range 1.69-4.28 kmol of MDEA/m3. Similarly eq 25 and 26 together with eq 27 and eq 16-19 can be solved to calculate the CO, solubility in aqueous BHEP solutions. A nonlinear equation solver routine, NDINVT, which is available at The University of British Columbia Computing Centre Library, was used to solve eq 13 and 14 and eq 25 and 26. This routine is based on the generalized secant method. The two unknowns to be determined are [H+] and a. Initial values of [H+] and a need to be provided for the iterative calculations, but they need not be very close to the final value. Typical initial values of [H+] and a are 10-6-10-8 and 1.5, respectively. No serious convergence problems were encountered.
+
Results and Discussion Tables I and I1 and Tables I11 and IV summarize the COBsolubility data for aqueous MDEA and BHEP solutions, respectively. The data are presented as a function of temperature, COPpartial pressure, and MDEA or BHEP concentration. The present method used to calculate the solubility is simple and somewhat approximate. To check the accuracy of the present method, experiments were also performed
E4 1
,
,
103
3
, , , , 3
57103
l,j
57104
CO2 PARTIAL PRESSURE ( k P a )
Figure 3. Present experimental and predicted (solid lines) COz solubilities in 1.69 M MDEA solutions as functions of COz partial pressure and temperature.
-
7
0-
10.'
10..
10-1 Iff
10'
Iff
Iff
10.
COZ PARTIAL PRESSURE ( k P a )
Figure 4. COz solubility data in 4.28 M MDEA solutions reported by Jou et al. (1982) and present model predictions (solid lines) as functions of COz partial pressure and temperature.
with a solution containing 4.28 kmol of MDEA/m3 at 100 "C and various CO, partial pressures. As seen from Figure 1,the present solubilities are slightly lower (average 8%) than those reported by Jou et al. (1982). An error analysis revealed that the small discrepancies arise from experimental and systematic errors. The estimated errors ranged from about 5% in the case of high CO, partial pressure to about 14% in the case of low C02 partial pressure. The present solubility data as well as those of Jou et al. (1982) are compared with the model correlations in Figures 2-5. Good agreement was found over wide ranges of temperature and C 0 2partial pressures. It was not possible to compare the present solubility data with the correlations based on the model of Jou et al. (1982). They did not give
2464
Ind. Eng. Chem. Res., Vol. 26, No. 12, 1987 0, 0
Table I. Experimental Solubility of COa in 4.28 M MDEA COz partial COz solubility, temp, OC pressure, kPa mol COz/mol MDEA 0.950 0.792 0.643 0.497 0.384 0.314 0.213 0.139 0.337 0.288 0.228 0.157 0.106 0.066 0.040 0.210 0.166 0.106 0.064 0.038 0.025 0.164 0.148 0.127 0.104 0.073 0.049 0.035 0.017 0.107 0.087 0.069 0.050 0.031 0.021 0.013
4930 3172 2103 1413 896 586 276 138 4516 362@ 2517 1413 724 310 138 4137 3034 1517 689 276 138 4550 3930 3240 2488 1517 793 414 138 3558 2733 2034 1344 621 345 172
100
140
160
180
200
o D E A 5 WTZ A B H E P 5 WT%
~1
$! , G
v
2
m w
1
,
,
~ ; l b ~, J
i
-1 30
0
z;
i-
c
-
0
IO'
3 5 71W
3
5 710.
3
5 710'
COZ PARTIAL PRESSURE ( k P a )
Figure 6. Solubility of C02 in 5 w t % DEA and 5 wt % BHEP solutions as a function of COz partial pressure at 100 "C. 0 ro
;z
0
1
A
o D E A 30 WTZ bDEA 25WTZ+BHEP 5 WTZ
,
1
N
1 ,
,
/ I , /
I
,
1
-
i
1
0
IW
3
3
5 7 1 0 .
5
710.
C O 2 PARTIAL PRESSURE ( k P a )
Figure 7. Solubility of COzin 30 wt % DEA and 25 wt % DFX plus 5 wt % BHEP solutions as a function of COz partial pressure at 40 "C.
c
4
-
10.8
10..
10-2
10'
101
IW
1w
IO'
COZ PARTIAL PRESSURE ( k P a )
Figure 5. COz solubility data in 2.0 M MDEA solutions reported by Jou et al. (1982) and present model predictions (solid lines) as functions of COz partial pressure and temperature.
A" 0 m
-
20; 0
10.
,
, , 3
, 571w
, , 3
571w
COZ PARTIAL PRESSURE ( k P a )
K1values at high temperatures, and it was difficult to obtain K , values from the logarithmic plots accurately. (Maddox et al. (1985) also presented some data on MDEA solubility, but they are not shown here due to the difficulty of reading their graphs accurately.) In order to compare the solubility of C02 in aqueous BHEP solutions with that in aqueous DEA solutions of similar concentrations, experiments were carried out with 5 w t % BHEP and 5 wt % DEA solutions at 100 "C. The results are present in Figure 6, and the solubility of C 0 2 in BHEP is seen to be less than that in DEA. Experiments were also carried out with 30 wt % DEA solutions and mixtures containing 25 wt % DEA and 5 wt % BHEP at 40 "C. These experiments were performed to establish whether DEA degradation and the formation
Figure 8. Present experimental and predicted (solid lines) COz solubilities in 0.287 M BHEP solutions as functions of COz partial pressure and temperature.
of BHEP reduce the C02 absorption capacity of the solution. The results are plotted in Figure 7. It is clear that formation of BHEP as a result of DEA degradation reduces the effective absorption capacity of the solution. The experimental BHEP solubility data are compared with the model predictions in Figures 8 and 9, respectively. As described in the Theory section, two sets of nonlinear equations have to be solved in order to calculate the C02 solubility in the C02-MDEA-H20 and C02-BHEP-H20 systems. The first equation (i.e., eq 13 or 25) relates the C 0 2 partial pressure to the acid gas loading and the hy-
Ind. Eng. Chem. Res., Vol. 26, No. 12, 1987 2465 Table 11. Experimental Solubility of COSin 1.69 M MDEA C02 partial COz solubility, temD. OC Dressure, kPa mol COJmol MDEA 1.304 100 4930 1.233 3930 1.120 2758 1.016 2000 0.904 1310 758 0.694 517 0.523 276 0.378 138 0.274 140 4309 1.003 2689 0.679 1827 0.484 1172 0.355 931 0.301 724 0.261 448 0.204 207 0.141 0.095 103 160 4275 0.710 2965 0.495 2068 0.380 1517 0.305 lo00 0.236 689 0.188 414 0.147 138 0.084 180 4137 0.504 2896 0.364 2068 0.283 1379 0.218 827 0.158 552 0.127 276 0.089 138 0.066 200 3585 0.349 2482 0.260 1793 0.213 1103 0.156 689 0.118 414 0.096
i 4
10.
3
5 7 1D
3
5 7 lo.
CO2 PARTIAL PRESSURE ( k P a )
Figure 9. Present experimental and predicted (solid lines) COz solubilities in 0.115 M BHEP solutions as functions of COz partial pressure and temperature.
drogen ion concentration, [H+], whereas the second equation (i.e., eq 14 or 26) gives the hydrogen ion concentration as a function of acid gas loading. If one could measure the hydrogen ion concentration directly and accurately, then the second equations would not be needed. Although pH electrodes capable of operating at high temperature and pressure are not readily available, they have been produced (Crolet and Bonis, 1983). Such electrodes might therefore provide a convenient way of determining acid gas loading, provided the Pco,, HCG,temperature, and the equilibrium constants are known. It appears worth-
Table 111. Experimental Solubility of COz in 0.2869605 M BHEP COSpartial COz solubility, temp, OC pressure, kPa mol COz/mol BHEP 40 3861 3.617 3137 3.007 2758 2.669 2137 2.116 1586 1.592 1137 1.178 621 0.640 276 0.292 138 0.168 100 3999 1.856 3551 1.664 2689 1.286 2034 0.986 1413 0.702 758 0.413 345 0.178 172 0.091 140 3654 1.235 2689 0.959 2379 0.841 1517 0.548 896 0.338 448 0.176 172 0.071 180 3999 1.064 3068 0.847 2620 0.710 1689 0.470 1000 0.273 517 0.148 193 0.063 Table IV. Experimental Solubility of COz in 0.115 M BHEP COz partial COPsolubility, temp, "C pressure, kPa mol COz/mol BHEP 40 4068 8.260 3447 6.992 2428 5.131 2206 4.599 1793 3.746 1379 2.841 lo00 2.124 483 1.122 172 0.470 3827 3.890 100 3447 3.490 3172 3.217 2758 2.851 2103 2.217 1379 1.475 896 0.987 586 0.686 207 0.271 140 3516 2.712 3103 2.474 2551 2.004 2137 1.688 1793 1.449 1379 1.074 896 0.746 483 0.423 207 0.191 180 3586 2.227 3206 2.034 2586 1.652 1620 1.072 655 0.489 241 0.123
while to explore this matter further in the future. Acknowledgment The financial assistance provided by the Natural Sciences and Engineering Research Council of Canada and
Ind. Eng. Chem. Res. 1987,26, 2466-2473
2466
Imperial Oil Ltd. is gratefully acknowledged. Nomenclature B = BHEP concentration, kmol/m3 BHEP = NJV-bis(hydroxyethy1)piperazine DEA = diethanolamine Hco = Henry's law constant, kPa/kmol/m3 K1-k5 = equilibrium constants Kbl = equilibrium constant for BHEP MDEA = methyldiethanolamine [MDEA] = total MDEA concentration, kmol/m3 MEA = monoethanolamine Pco, = C 0 2 partial pressure, kPa R = CPHdOH T = temperature, K Greek Symbols LY
= C 0 2 loading, mol of C02/mol of solvent
[ ] = molar concentration of species, kmol/m3 Registry No. MDEA, 105-59-9;BHEP,122-96-3;COz, 124-
38-9.
Blanc, C.; Elgue, J.; Lallemand, F. J. Hydrocarbon Process. 1981, 60(8),111. Chakma, A.; Meisen, A. Proceedings of the 35th Canadian Chemical Engineering Conference, Calgary, Alberta, 1985, Vol. 1, p 37. Crolet, J. L.; Bonis, M. R. Corrosion 1983, 39(2), 39. Deshmukh, R. D.; Mather, A. E. Chem. Eng. Sci. 1982,36(2), 365. Frazier, H. D.; Kohl, A. L. Znd. Eng. Chem. 1950, 42(11), 2288. Guggenheim, E. A. Phil. Mag. 1935,19, 588. Jou, F.; Mather, A. E.; Otto, F. D. Znd. Eng. Chem. Process Des. Deu. 1982, 21, 539. Kennard, M. L.; Meisen, A. Znd. Eng. Chem. Fundam. 1985,24,129. Kent, R. L.; Eisenberg, B. Hydrocarbon Process. 1976, 55(2), 87. Kohl, A. L.; Riesenfeld, F. C. Gas Purification, 4th ed.; Gulf Publishing: Houston, 1985. Maddox, R. N.; Bhairi, A.; Mains, G. J.; Shariat, A. Acid and Sour Gas Treating Processes; Newman, S. A., Ed.; Gulf Publishing: Houston, 1985; p 212. Peng, D. Y.; Robinson, D. B. Znd. Eng. Chem. Fundam. 1976,15(1), 59. Polderman, L. D.; Steele, A. B. Oil Gas J. 1956, 54(65), 206. Reid, R. C.; Prausnitz, J. M.; Sherwood, T. K. The Properties of Gases and Liquids, 3rd ed.; McGraw Hill: New York, 1977. Smith, R. F.; Younger, A. H. Hydrocarbon Process. 1972,51(7), 98. Vidaurri, F. C.; Kahre, L. C. Hydrocarbon Process. 1977,56(11),333.
Literature Cited
Received for review October 6, 1986 Revised manuscript received July 15, 1987 Accepted August 9, 1987
Benedict, M.; Webb, G . B.; Rubin, L. C. Chem. Eng. Prog. 1951,47,
419.
A Comparative Study of Catalyst Deactivation in Integrated Two-Stage Direct Coal Liquefaction Processes Frances V. Stohl* and Howard P. Stephens Sandia National Laboratories, Organization 6254, Albuquerque, New Mexico 87185
Catalyst samples from Wilsonville runs 242, 246,247, and 248 have been characterized and tested for hydrodesulfurization activity in order to determine the effects of different process configurations and coals on the causes and rates of deactivation. Initial rapid decreases in activity were caused by the accumulation of carbonaceous deposits within the catalyst. Variations in the amount of deposition were due to the process configuration but not the coal type. The process configuration with the most preasphaltenes and unconverted coal in the hydrotreater feed yielded the greatest buildup of carbonaceous deposits and the most deactivation. After the initial high losses, the activities continued to decrease with catalyst age due to the accumulation of contaminant metals in the catalyst. The rates of contaminant metals buildup varied significantly from run to run. Deactivation by carbonaceous deposits and contaminant metals was due to both poisoning of active sites and decreases in effective diffusivities. In addition, sintering of the active metals was observed in the aged catalysts.
I. Introduction During the past several years, direct coal liquefaction processes have evolved from single-stage processes to two-stage processes (Neuworth and Moroni, 1984; Schindler et al., 1982; Rao et al., 1983). Dissolution of the coal occurs in the first stage, and catalytic upgrading takes place in the second stage. Separation of the two stages enables each to be operated under suitable conditions for the reactions that occur in that stage. Current two-stage processes are operated in the integrated mode in which a portion of the second-stage product is recycled back to the first stage as a solvent to be mixed with the coal. Researchers at the Wilsonville Advanced Coal Liquefaction R & D Facility have evaluated the use of three integrated two-stage process configurations for direct liquefaction of two coals (Lamb et al., 1985; Moniz and Nalitham, 1985; Moniz et al., 1983, 1984; Gough et al., 1985). Results showed that resid (material that is nondistillable at 316 "C and 0.1 mmHg) conversion, which was used as an indicator of catalyst performance, decreased 0888-5885/87/2626-2466$01.50/0
with time for all process configurations and coals. However, the rates of decrease in resid conversion varied among runs (Moniz and Nalitham, 1985). The greatest rates of decrease in resid conversion were early in all the runs. Most studies of aged coal liquefaction catalysts have been of catalysts derived from single-stage or nonintegrated two-stage processes (Thakur and Thomas, 1984; Cable et al., 1981, 1985; Stohl et al., 1987). There have only been a few studies of catalysts from integrated two-stage processes (Stiegel et al., 1985; Freeman et al., 1985). The objective of our work was to determine the causes for the differences in catalyst deactivation in the Wilsonville runs. To accomplish our objective, we characterized and tested catalyst samples from four Wilsonville liquefaction runs: 242, 246, 247, and 248. Catalyst deactivation (Ocampo et al., 1978; Kovach et al., 1978; Furimsky, 1979; Thakur and Thomas, 1984) is caused by the accumulation of carbonaceous deposits and contaminant metals on coal liquefaction catalysts. The buildup of carbonaceous deposits is very rapid, whereas 0 1987 American Chemical Society
