I n d . E n g . Chem. Res. 1988,27, 324-328
324
Xco2,3i= weight fraction of C02as CaCO, in the retorted solids Xco2,20= weight fraction of C02as MgCO, in the spent solids Xco2,30= weight fraction of COPas CaC03in the spent solids XH = weight fraction of H in carbon residue
Xi= weight fraction of particles in class i
X s = weight fraction of S in retorted solids 2 = distance from air distributor Greek Symbols h p b = bed pressure drop h p d = distributor pressure t = void fraction t b = bubble fraction
drop
emf = void fraction at minimum fluidization conditions p = gas velocity pg = density of gas ps = particle density
4 = particle sphericity
Literature Cited Baughman, G. L. Synthetic Fuels Data Handbook; Cameron Engineers: New York, 1978; p 14. Braun, R. L. Report UCRL-53119; 1981; Lawrence Livermore Laboratory, London. Campbell, J. H. 1978 Report UCRL-52089, Part 11, 1978; Lawrence Livermore National Laboratory; London.
Cranfield, R. R.; Geldart, D. Chem. Eng. J . 1972, 3, 211-231. Fluidization Engineering; Robert E. Krieger: Huntington, NY, 1977. Geldart, D. EPRI CS-2094, Oct 1981; EPRI, Washington, DC. Hall, R. N., AIChE Meeting, Anaheim, CA, June 1982. International Critical Tables; McGraw-Hill: New York, 1930; Vol. VII, p 292. Knowlton, T. M., AIChE Meeting, New York, Dec 1-5, 1974. Mori, S.; Wen, C. Y. AIChE J. 1975, 2I(1), 109-115. Patzer, J. F.; Carr, N. L., Presented at the AIChE National Meeting, Anaheim, CA, 1984. Rammler, R. W., Oil Shale Processing Technology; Dean Allred, V., ed.; Ellenar Graphics: New York, 1982. Sohn, H. Y.; Kim, S. K. Znd. Eng. Chem. Process Des. Deu. 1980,19, 550. T a " , P. W. "Combustion of Pyrolyzed Carbon Containing Solids in Staged Turbulent Bed". U S . Patent 4 336 128, 1982. Ta", P. W.; Bertelsen C. A.; Handel, G. M.; Spars, B. G.; Wallman, P. H. Energy Prog. 1982,2(1), 37. Vasalos, I. A.; Tatterson, D. F.; Furlong, M. W.; Kowalski, T. L.; So, B. Y. C., Annual AIChE Meeting, San Francisco, 1984. Wallman, P. H.; Tamm, P. W.; Spars, B. G. Prepr.-Am. Chem. SOC., Diu. Pet. Chem. 1980, 25(3), 70. Wen, C. Y.; Yu, Y. H. AIChE J . 1966,12, 610.
Received for review February 21, 1986 Revised manuscript received May 27, 1987 Accepted October 8, 1987
SEPARATIONS Absorption of Carbon Monoxide into Aqueous Solutions of K2C03, Methyldiethanolamine, and Diethylethanolamine C. J. Kim,* Alan M. Palmer, and George E. Milliman E x x o n Research and Engineering Company, Corporate Research Laboratories, Annadale, N e w Jersey 08801
Rates of CO absorption into 1 M aqueous solutions of K2C03,diethylethanolamine (DEAE), and methyldiethanolamine (MDEA) were measured in a stirred tank reactor a t 348-398 K and a t CO pressures in the range of 7.5-31 bar. The absorption behavior showed the absence of any mass-transfer limitations, and detailed kinetic analyses established t h a t a reaction step, CO + OHHCOO-, is rigorously controlling the absorption rates in all systems examined in this study. Comparative discussions on the kinetics of C 0 2 OH- vs CO OH- are given.
-
+
Removal of C02and H2Sfrom gases is usually achieved by regenerative absorption into -aqueous solutions containing bases such as amine-promoted potassium carbonate solutions and aqueous amines. In cases when the feed gas contains carbon monoxide, the gas-treating solutions become gradually deactivated due to the formation of stable formate salts. A brief account of such deactivation of diethanolamine-promoted potassium carbonate solutions in commercial plants was described by Eickmeyer (1962). The reaction of CO with K2C03has been shown by Yoneda et al. (1943) to occur by CO + K2C03+ H 2 0 KOOCH + KHC03 (1) According to the thermodynamic data compiled by Latimer (1952),the standard free energy of reaction 1is -19.2 kJ/mol, while that of reaction 2 is -14.4 kJ/mol. Under
-
*Present address: 1240 Rattlesnake Bridge Road, Bedminster, N J 07921. 0888-5885/88/2627-0324$01.50/0
+
COZ
+ K&03 + H2O
+
2KHCO3
(2)
the usual gas-treating conditions, regenerative CO desorption by the reverse process of reaction l cannot be expected because the facile COz desorption via the reverse step of reaction 2 depletes HC03-, the proton-donating species needed for decomposition of the formate ion. It is thus expected that CO absorption leads t o degradation of the basic components of the gas-treating solutions into nonregenerable formate salts. The extent and severity of such degradation depend on the rates of CO absorption, and the present study was conducted to define the basic chemistry involved in the reactions of CO with K,CO,, methyldiethanolamine (MDEA), and diethylethanolamine (DEAE) in aqueous solutions. I
I.
Experimental Section Absorption Rate Measurement Procedure. A 300cm3 stirred autoclave (Autoclave Engineers Hastloy-C 0 1988 American Chemical Society
Ind. Eng. Chem. Res., Vol. 27, No. 2, 1988 325 autoclave with a magnedrive assembly) was equipped with a precision pressure gauge, a heating jacket controlled by a proportionating temperature controller, liquid and gas sampling lines, and a thermocouple inserted into the liquid layer. A 150-cm3charge of 1M solution of K2C03,MDEA, or DEAE was introduced, and the vapor space was purged with helium. The reactor was then pressurized with helium to about 7 bar, heated to a preset temperature, and after a steady temperature control ( f l K) was attained, the pressure was released to 0 gauge pressure. CO was then introduced to the desired pressure level to start the reaction a t a stirring rate of 1000 rpm. The rate of CO uptake was measured by recording the cumulative pressure drop with time, while the total pressure was allowed to fluctuate &lo% of the set level during a run. The effective CO partial pressure was calculated from the average total pressure and the steam pressure. The amounts of absorbed CO calculated from pressure measurements were checked against liquid sample titration results, and the agreements were within f5%. Liquid Sample Analysis Procedure. Control experiments using synthetic solutions containing known amounts of K2C03,KHC03, NaOOCH, and/or an amine showed that a complete solution characterization can be carried by the following combination of procedures. The concentrations of carbonate and bicarbonate species were determined by titration with a standard hydrochloric acid solution employing Metrohm E-536 automatic recording titrator. At the first equivalence point, the reaction, CO?- + H+ HCO,, is complete, and a t the second equivalence point, the second neutralization reaction, HC03- + H+ H 2 0 + C02, is complete. The amounts of CO2- and HC03- in a sample can thus be determined from the titration result. The amount of unreacted amine in solution can also be defined by a similar procedure. The formate salt concentration was measured by a separate procedure using an ion chromatograph (Dionex Model 16). A 0.001 25 M NazB40, solution was used as the eluent on anion separator and supressor columns. Another set of analyses was carried out to examine whether the spent amine solutions contained degradation products such as formyl esters of the alkanolamines. Accordingly, samples of amine solutions were saturated with K2C03, extracted with isopropyl alcohol, and analyzed on a Tenax GC column using an FID detector. No detectable peaks attributable to degradation products were observed in the MDEA or DEAE samples.
-
-
Absorption of CO into K&03 Solutions The rate of CO absorption can be expressed by time dependent variation of a which is defined as a = moles of CO absorbed/ moles of K2C03initially present (3) In this study, a is pegged to the amount of formate ion present in solution and does not include dissolved CO. On the basis of the solubility data in pure water (Seidell, 1958), the concentration of physically dissolved CO is estimated to be very low under the present measurement conditions, below about 0.007 M, and also expected to remain more or less invarient for the duration of a given run. Thus, the contribution by physically dissolved CO to the differential amount of CO absorbed becomes negligible and a meaningful kinetic analysis can be carried out using CO partial pressure in place of the term representing physically dissolved CO. The rates of CO absorption into 1 M K2C03solutions were measured a t 348-398 K and at CO partial pressures of 7.5-31.6 bar. The plots in Figures 1and 2 illustrate the
0'5r----
01 0
I
100
I
1
200
300 Reaction Time, Min
1
400
)O
Figure 1. Rates of CO absorption into 1 M KzCO3 solution at 373 K (A) Pco = 31.6 bar, (0)Pco = 18.1 bar, (v)Pc0 = 7.5 bar. 0.71
CI
1
I
I!
100
200
300
400
500
$00
Reaction Time, Min
Figure 2. Rates of CO absorption into 1 M KzC03 solution at Pco = 18.1 bar: (A)398 K, (0) 373 K, (V)348 K.
dependencies of the rate on the CO partial pressure and temperature. In each run, the solution was analyzed to confirm the formation of equimolar mixture of KOOCH and KHC03 in amounts corresponding to the stoichiometry of eq 1. In some cases, however, the amount of KHC03was slightly lower than that of KOOCH, indicating that partial C02 desorption through the reverse step of reaction 2 was taking place during measurements. Under the typical measurement conditions of CY < 0.4 in a closed system, the amount of COz distributed in the gas phase was estimated to be less than 5% of the C02 equivalent in solution (one-half the amount of KHC03 in solution). Data analysis was thus carried out neglecting the contributions of reaction 2. In 1960 Kodama et al. (1960) found that the rate of CO absorption into aqueous NaOH solutions is strictly controlled by chemical rate at 308-358 K and follows the rate expression d[HCOO-]/dt = KPco[OH-] (4) This equation represents a second-order expression in which the activity of dissolved CO in solution is substituted for by a pressure term. The reaction of CO with K2C03 was assumed to occur, as proposed by Royen and Erhard (19561, according to the reactions steps C032- + H 2 0 HC03- + OH(5)
OH-
--
+ CO
HCOO-
(6)
326 Ind. Eng. Chem. Res., Vol. 27, No. 2, 1988
The rate of overall reaction may then be expressed by 0'15----
where K , is the equilibrium constant of reaction 5 expressed as [OH-][HC03-]/[C0,2-]. As the species concentrations may be expressed as [C02-] = Co(1 - a) [HC03-] = [HCOO-] = COa where Co is the initial K2C03concentration (in M), eq 7 becomes
-. 100
Integrating and substituting Co = 1, we obtain -[ln (1 - a) - a] = ( k K g c o ) t
(9)
Thus, if the reaction sequence represented by eq 5 and 6 is valid, linear correlations should be found when values corresponding to -[In (1- a)- a ] are plotted against time. Indeed, satisfactory linear plots are observed as shown in Figures 3 and 4. The data at 373 K further show that the slopes of the correlation lines are a linear function of CO partial pressure, which is consistent with the expression in eq 9. The magnitudes of the term, kK,, were then calculated from the slopes and summarized in Table I. The temperature dependency of kK, is represented by activation energy of 123 kJ/mol which shows that the rate of CO absorption into KzC03 solutions is very sensitive to temperature variations and also that mass-transfer limitations are not important a t present measurement conditions. Further analysis to decouple the kK, term was carried out as follows. The thermodynamic K5 value a t infinite dilution conditions, K5",is given by K5" = Kw"/KII"
K5"
Yco32YHCO~-YOH-
The activity coefficients of ionic species are calculable according to the extended Debye-Huckel equation (Davies, 1938; Perrin et al., 1981): -(log
+
= ~ ~ ~ 2 1 0 , 5 / (1 01. 5 ) - O.IZ:I
(12)
where yi is the activity coefficient of species i, Ziis the charge number of species i, A is the Debye-Huckel parameter, and I is the ionic strength defined by
I = 1/2cz:ci
(13)
with Cj being the concentration of species i. This equation is useful for solutions of low ionic strengths, and the calculated K6/K5" values decrease with increasing ionic strength as 0.78 ( I = 0.01) and 0.53 ( I = 0.1) a t 373 K (Table 11). This trend suggests that the K5/K5"value for
400
Figure 3. Correlations of -[ln (1- a) - a]vs time for the CO-K2C03 absorption data at 373 K (A)Pco = 31.6 bar, ( 0 )PCO= 18.1 bar, (v)Pco = 7.5 bar. 0.31
I
(10)
where Kwois the HzO dissociation constant and Kn" is the second ionization constant of carbonic acid. As the values of K," and Kno at 348-398 K are available in the literature (Perrin, 1982)) K," can be readily calculated. The K5 values, expressed in concentration terms a t nonideal solution conditions, are then related to K5" by the activity coefficient term
-K5- -
200 300 Time, Min
10 Time, Min
Figure 4. Correlations of -[ln (1- a ) - a]vs time for the CO-K2C03 absorption data at Pco = 18.1 bar: (A)398 K, ( 0 )373 K, (v)348 K. Table I. Kinetic Constants for the CO-K,CO, Reaction slope," kK5,*min-' temp, K CO pressure, bar min-' bar-' M-' 398 17.3 2.2 x 10-3 1.27 x 10-4 373 31.6 2.8 X 8.8 X lo4 18.1 1.45 x 10-4 8.0 x 104 7.5 5.9 x 10-5 7.9 x 104 8.2 X lo4 av 1.09 x 10-5 6.0 x 10-7 348 18.1 "From -[ln (1 - a ) - a] vs time plots in Figures 3 and 4. *For definitions, see eq 5-9 in text.
the working solution, 1 M K&O, solution ( I = 3.0),is considerably lower than unity, while it is noted that eq 12 may not be applied to solutions of high ionic strength. Accordingly,an alternative approach to determine K5 was adopted. The k values reported by Kodama et al. (1960) for CO + OH- reactions a t 318-343 K were defined for 3 M NaOH solutions ( I = 3.0). We reasoned that Kodama's k values should represent the present CO-K2C03 system fairly accurately because the ionic strengths of the solutions are closely matched and also because the k value was derived from the CO partial pressure term which, in the absence
Ind. Eng. Chem. Res., Vol. 27, No. 2, 1988 327 Table 11. Estimated Equilibrium and Kinetic Constants Related to Reaction of CO in 1 M K&03 Solution a t 348 K a t 373 K a t 398 K kK5 ( I = 2.8); m i d bar-' 6.0 X 8.2 X lo4 1.27 X lo4 M-1 k ( I = 3.0),*min-' bar-' 6.9 X 5.3 X 3.1 X K5 ( I = 2.8), M 8.7 x 10-4 1.6 x 10-3 4.1 x 10-3 K5O ( I = O),' M 2.82 X 7.76 X 2.09 X lo-* K5/K5' ( I = 2.8) 0.31 0.21 0.20 K 5 / K 5( I = O . l ) d 0.56 0.53 0.51 K5/K50( I = O.O1)d 0.79 0.78 0.76 'See Table I. *Extrapolated from Kodama's data (1960). 'Calculated according to eq 10 using data compiled by Perrin (1965). dCalculated by using eq 12. 0.4
I
1
0.01
1/ 00 Time, Min
a
Figure 6. Correlations of -[ln (1- a) - a]vs time for the CO-amine absorption data a t 373 K: (A) DEAE, (0) MDEA.
a CO + OH- route, the following simple reaction sequence was set up by treating amines as simple bases: O?
0
I
500
I
I
1000 1500 Reaction Time, Min
I
AM
+ H2O & AMH+ + OH7
2000
CO
+ OH-
HCOO-
(14) (15)
Figure 5. Rates of CO absorption into 1 M amine solutions at 373 K: (A) DEAE a t Pc0 = 18.1 bar, (0) MDEA a t Pco = 19.3 bar.
of mass-transfer limitations, should be a direct measure of the activity of CO dissolved in solutions. Thus, k values at 348-398 K were calculated by extrapolating Kodama's data and subsequently used to calculate K5 from the previously established kK5values. The results summarized in Table I1 show that the magnitude of K5/K50 a t I = 2.8 is 0.21 (at 373 K), which is consistent with the previously calculated trend at I = 0.1 and 0.01. In fact, the calculated magnitudes of K5/K50 in this study are very similar to those reported by Nasanen (1946), who reported K5/K50 in solutions containing controlled amounts of KCl decreased with increasing ionic strength as 0.60 (I = O.l), 0.37 (I = 0.5), 0.33 ( I = LO), and 0.27 ( I = 2.0). It thus appears that the kinetic analysis results listed in Table I1 are entirely consistent with a view that CO absorption into K2C03solutions occurs via reaction steps represented by eq 5 and 6.
These equations parallel eq 5-7 for the CO-K2C03 system, and the same kinetic analysis procedure may be applied by broadening the definition of a as a = moles of CO absorbed/ moles of K2C03or amine initially present
Absorption of CO into DEAE and MDEA Solutions
where KAo is the thermodynamic dissociation constant of protonated amine species a t infinite dilution conditions. KAo values of DEAE and MDEA in the form of pK, [= -(log KAo)]are available in the temperature range 293-333 K (Perrin, 1965), and the extrapolated pK, values at 373 K are 8.4 (DEAE) and 7.4 (MDEA). As the ionic strength of 1 M amine solutions in the measurement range of a = 0 . 2 is low ( I = 0.1 average), the activity coefficient term, Y ~ + Y o H - was , calculated by using eq 11as 0.59 (a= 0.06), 0.53 (a= 0.10), and 0.48 (a= 0.16). By use of Y ~ H + Y O H = 0.53 and -(log Kwo)= 12.26 at 373 K, the equilibrium M constant K14 a t 373 K was tabulated as 2.9 X M (MDEA). The results of this (DEAE) and 2.3 X thermodynamic approach to define K14 may now be used to extract k from the previously determined kK14 values. min-' bar-' (from The calculated k values are 5.4 X min-l bar-l (from the the DEAE data) and 6.5 X MDEA data), which are in good agreement with lz = 5.3
The primary objective of measuring rates of CO absorption into aqueous tertiary amine solutions was to examine whether the underlying chemistry for CO-amine reactions is the same as that described for the CO-K2C03 system. As the ionic strengths of 1M amine solutions at low conversion conditions are sufficiently low for meaningful application of the extended Debye-Huckel equation, the amine system is expected to provide a means to perform kinetic analysis using thermodynamic data available in the literature. DEAE and MDEA were chosen for this study largely because their pK, values are available for a wide range of temperatures from the compilation of Perrin (1965). The rates of CO absorption are illustrated in Figure 5. Assuming that the CO-amine reaction also occurs via
The plots of -[ln (1 - a) - a] vs time gave good linear correlations as shown in Figure 6, and the kKI4values a t 373 K were extracted from the slopes as 1.54 X lo4 (DEAE) and 1.50 X lo-' min-I bar-' M-l (MDEA). The kinetic behavior of the CO-amine reaction is thus consistent with the formulation in eq 14-16, and this is further substantiated by the discussion given below. The equilibrium constant K14 is accessible by v o
328 Ind. Eng. Chem. Res., Vol. 27, No. 2, 1988
x min-' bar-' previously derived from Kodama's data (1960). These results together with the previous results for the CO-K2C03 system thus provide strong evidence supporting the view that the rate of CO absorption into aqueous solutions containing K,C03, DEAE, MDEA, or NaOH is controlled by its reaction with OH-. The remarkably consistent k value determined for solutions of widely different solution properties also supports the point that the activity of CO involved in the reaction is well represented by the CO partial pressure term. Discussion The rates of CO absorption into solutions containing K2C03, DEAE, and MDEA are much slower than the corresponding rates of C02 absorption. The results of kinetic analysis seem to establish that CO absorption is free of any mass-transfer limitations even a t moderately high temperatures up to 398 K but is controlled by a chemical rate representing a step, CO + OH- HCOO-. Thus, the practical problem of assessing the CO-induced degradation of gas-treating solutions can be quantified by defining three parameters, the chemical potential of CO (partial pressure of CO), the activity of OH-, and temperature. For systems in which a OH-activity scale can be defined, the prediction on the expected rate of degradation due to CO absorption becomes a fairly simple task. Conversely, as noted before, a series of careful CO absorption rate measurements may provide valuable information on the OH- activity level of various gas-treating solutions at conditions of practical interest. It is, however, noted that the most important parameter that controls the CO absorption rate into buffered gas-treating solutions is the temperature, as exemplified by the high apparent activation energy of 123 kJ/mol for the CO-K2C03system. Since the C02 or H2S absorption rates are much less sensitive to temperature changes, the CO-induced degradation problem is expected to become progressively more severe when the process temperature becomes higher. Finally, it is of interest to compare the reactivity of OHtoward CO and C02. Pinsent et al. (1956) studied the intrinsic kinetics of the C 0 2 0 + OH- reaction by employing a rapid mix technique. The second-order rate constant of at 323 K, which this reaction, kcot, is 4.8 X lo5 M-' shows that the reaction is extremely facile; e.g., C 0 2 in a solution containing M concentration of OH- would react with a half-life of approximately 1.4 X s. The relatively low activation energy (55.4 kJ/mol) is also consistent with the noted facility of the reaction which may be depicted by the simple nucleophilic addition reaction shown below:
-
0
HO-
+ CII II
0
- Ho-i 0
The reaction of CO with OH-, on the other hand, is sluggish a t 323 K. A comparative second-order rate constant for the CO[,, + OH- reaction, kco, is not obtainable a t this time, but an approximate value may be estimated using the solubility of CO in pure water (Seidell, 1958): 0.7 X M at Pco = 1bar and 323 K. If k = 6.64 X bar-l min-' a t 323 K is combined as defined before, the magnitude of kco a t 323 K is of the order of 2 X M-' s-l. This number shows that the rate of CO reaction is slower than that of C02 by a factor of lo8, while the ac-
tivation energy of the CO reaction is 92 kJ/mol, which is significantly higher when compared with 55.4 kJ/mol for the CO, reaction. These characteristic kinetic parameters suggest that the CO + OH- reaction involves a high-energy intermediate (or an activated transition complex). A possible reaction route is given below:
The precise nature of the CO + OH- reaction, however, remains unknown a t this time, and a comprehensive understanding of the formate ion-CO chemistry is needed, especially in light of the recent interest in the role of formate ion as a homogeneous catalyst for water gas shift reaction at high temperature/pressure conditions (Elliott et al., 1983; Zielke et al., 1976).
Nomenclature A = Debye-Huckel parameter Co= initial concentration of KzC03or amine, M I = ionic strength as defined in eq 13 k = second-order rate constant for CO + OH- as defined in eq 4, min-' bar-' kco, kco = second-order rate constants for COz + OH- or CO + OJ$ in solution, s-1 M-1 K = equilibrium constants defined by concentration terms KO = equilibrium constants defined by activity terms P = pressure t = time 2 = charge number Greek Symbols a = moles of CO absorbed per mole of KzC03or amine initially present in solution y = activity coefficient Registry NO.DEAE, 100-37-8; MDEA, 105-59-9; CO, 630-08-0; KzC03, 584-08-7.
Literature Cited Davies, C. W. J. Chem. SOC.1938, 2093. Eickmeyer, D. C. Chem. Eng. Prog. 1962,58, 89. Elliott, D. C.; Hallen, R. T.; Sealock, L. J. Ind. Eng. Chem. Prod. Res. Dev. 1983, 22, 431. Kodama, S.; Tomihisa, N.; Shimamura, I,; Fukui, K. Kogyo Kagaku Zasshi 1960, 63, 1733. Latimer, W. M. Oxidation Potential, 2nd ed.; Prentice-Hall: Englewood Cliffs, NJ, 1952. Nasanen, R. Acta Chim. Fenn. 1946, Bl9, 90. Perrin, D. D. Dissociation Constants of Organic Bases in Aqueous Solutions; Butterworths: London, 1965; Supplement 1972. Perrin, D. D. Ionization Constants of Inorganic Acids and Bases in Aqueous Solutions, 2nd ed.; Pergamon: New York, 1982. Perrin, D. D.; Dempsey, B.; Serjeant, E. P. pKa Predictions for Organic Acids and Bases; Chapman and Hall: London, 1981. Pinsent, B. R. W.; Pearson, W. L.; Roughton, F. J. W. Trans. Faraday SOC.1956,52, 1512. Royen, V. P.; Erhard, F. Erdol Khole 1956, 9, 19. Seidell, A. Solubilities of Inorganic and Metal Organic Compounds; McGregor-Werner: Washington, D.C., 1958. Yoneda, K.; Honda, Y.; Momiyama, N.; Abe, R. J . SOC.Chem. Ind. Jpn. 1943, 46, 554. Zielke, C. W.; Rosenhoover, W. A,; Gorin, E. Am. Chem. SOC.Fuels Prepr. 1976, 21, 163.
Received for review March 31, 1986 Revised manuscript received September 17, 1987 Accepted October 8, 1987
