Ind. Eng. Chem. Res. 1987,26, 27-31 Bird, R. B.; Stewart, W. E.; Lightfoot, E. N. Transport Phenomena; Wiley: New York, 1960. Bomben, J. L. Ph.D. Thesis, University of California,Berkeley, 1981. English, A. C.; Dole, M. J. Am. Chem. SOC.1950, 72, 3261-3267. Flemings, M. C. Solidification Processing; McGraw-Hill: New York, 1974. Frank, F. C. Proc. R. SOC.London, Ser. A 1950, 20A, 586-599. Gosting, L. G.; Morris, M. S. J.Am. Chem. SOC.1949, 71, 1998-2006. International Critical Tables; McGraw-Hik New York, 1928; p 263. Mullins, W. W.; Sekerka, R. F. J . Appl. Phys. 1963, 34, 323-329. Newman, J. Ind. Eng. Chem. Fundam. 1968, 7, 514-517.
27
Newman, J. S. Electrochemical Systems; Prentice-Hall: Englewood Cliffs, NJ, 1973. Scriven, L. E. Chem. Eng. Sci. 1959, 10, 1-13. Sekerka, R. F. J. Appl. Phys. 1965,36, 264-268. White, R.; Mohr, C. M.; Fedkiw, P.; Newman, J. The Fluid Motion Generated by a Rotating Disk: A Comparison of Solution Techniques; University of California: Berkeley, 1975; LBL-3910. Received for review April 4, 1983 Revised manuscript received January 13, 1986 Accepted March 12, 1986
Reactions of Carbon Dioxide and Hydrogen Sulfide with Some Alkanolamines Robert N. Maddox,*t Gilbert J. Mains,$ and Mahmud A. Rahmant Department of Chemistry and School of Chemical Engineering, Oklahoma State University, Stillwater, Oklahoma 74078
Carbon dioxide and hydrogen sulfide were interacted with anhydrous diglycolamine (DGA), di-2propanolamine (DIPA), methyldiethanolamine (MDEA), and dimethylethanolamine (DMEA) and the reaction products dissolved in DCC13 for NMR analysis. The resultant NMR spectra are compared with the NMR spectra of the unreacted amines at similar concentrations. -Evidence for a Lewis acid-base complex involving H2Sand DIPA was found. NMR absorptions attributable to the formation of protonated amine:carbamate salts were found for both DGA and DIPA. Although no NMR evidence was obtained for reaction products from DGA, MDEA, or DMEA using either H2S or C02, pressure measurements suggest that Lewis acid-base adducts are reversibly formed in these systems and decompose prior to NMR analysis. The removal of carbon dioxide and hydrogen sulfide from gaseous mixtures is of vital importance in the natural gas and petroleum gas processing industry. In principle, many methods are available for the removal of acid gases, but the use of liquid solvents is economically attractive (Maddox, 1982). The amine solvents which are of commercial interest and are considered in this study are /3,/3'-hydroxyaminoethyl ether (diglycolamine, DGA), di2-propanolamine (DIF'A), methyldiethanolamine (MDEA), and dimethylethanolamine (DMEA). The reactions between carbon dioxide and amines have been extensively studied, but the kinetics and relevant products are still a matter of speculation and dispute (Astarita et al., 1982). Similar uncertainty exists in the case of hydrogen sulfide. The objective of this study is to ascertain the reaction products between the acid gases and anhydrous alkanolamines. The reaction products were analyzed by means of a 100-MHz nuclear magnetic resonance spectrometer.
Reactions of Hydrogen Sulfide with Amines When H2S comes in contact with an aqueous amine, or any alkaline solution, it dissociates into the bisulfide ion, SH-, by a proton-transfer reaction: H2S + amine (H:amine)+ + SH(1)
-
This may be considered to be very rapid when compared to the rate of diffusional processes (Maddox, 1982; Astarita et al., 1983). The bisulfide ion does not dissociate further into the sulfide ion, S2-,except in strong hydroxide solutions (Ladda and Danckwerts, 1981). If the amine is anhydrous, reaction 1 results in an acid-base complex,
* Author
*
to whom all correspondence should be addressed. School of Chemical Engineering. Department of Chemistry. 0888-5885/87/2626-0027$01.50/0
amine:H2S,which is not dissociated into ions in the low dielectric amine solvents.
Reactions of Carbon Dioxide with Amines The reactions of carbon dioxide with aqueous amines may result either in the formation of carbonic acid (H2C03) and subsequent, rapid protonation of the amine by a reaction analogous to reaction 1for HzS or in the formation of a carbamate salt depending upon whether the amine has a N-H bond. Thus, primary (RNH2) and secondary amines (RzNH)can undergo carbamate formation, whereas a tertiary amine (R,N) cannot. The formation of the carbamate is believed to occur in two steps, the rate-determining step is the insertion of C02 into the N-H bond which probably involves the reversible formation of a Lewis acid-base adduct COZ + R2NH R2HN:CO, RZNCOOH (2) followed by the very rapid protonation of another amine molecule. R2NH + HOOCNRi [R2NH2+][-00CNRJ (3) In water solutions the carbamate salt from reaction 3 is dissociated into free ions. However, in anhydrous amines the salt remains as associated ion pairs because of the low dielectric constants of these solvents. The kinetics of the reaction have been shown to be second order, consistent with reaction (Sharma and Danckwerts, 1963; Hikita et al., 1979). Of course at C02/amine loadings greater than 112, reaction 2 cannot occur. Initially, reaction 2 is so rapid that mass transfer into and away from a thin reaction zone at the surface of the solution complicates the interpretation of many kinetic measurements (Danckwerts, 1970). The formation of carbonic acid can occur only in aqueous solutions (Hikita et al., 1977). The hydration reaction 4 (Astarita et al., 1983; Sigmund et al., 1981) is known to have a half reaction time of about 23 s at 25 O C . -+
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28 Ind. Eng. Chem. Res. Vol. 26, No. 1, 1987
COz + H2O HzC03 (4) Indeed, the equilibrium constant for reaction 4 is 2.57 X loT3,and the dissolved carbon dioxide is primarily in the form of COz rather than carbonic acid at 25 "C. The subsequent dissociation of carbonic acid into a proton and bicarbonate ion, K1 = 1.70 X at 25 "C, HzC03 H+ + HC03(5) -+
lo-" -
and dissociation of bicarbonate into a proton and carbonate ion, K , = 4.69 X g-mol/L a t 25 OC, HC03- H+ + C032(6) achieve equilibrium very quickly. The protons react rapidly with amine (H:amine)+ (7) H+ + amine to form tertiary amine ions, reaction 7, the equilibrium constant of which depends upon the particular amine but varies roughly between lo4 and 1O'O. Despite the speed of reactions 5 and 6, a t COz/amine loadings of less than 0.5, these reactions are of little practical importance for primary and secondary amines since reaction 2 is much faster than reactions reported by Sharma and Danckwerts (1963) and Hikita et al. (1977). Reaction 7 remains important, however, since dissociation of the carbamic acid formed by reaction 2 into protons and carbamate ions and the subsequent removal of the protons by reaction 7 is almost certainly the primary mechanism by which quaternary amine ions are formed in aqueous systems. In addition, alkanolamines have the hydroxyl functional group, which may react with carbon dioxide in a manner similar to the hydration reaction ROH C 0 2 ROCOOH (8) to form esters of carbonic acid (reaction 8). These esters are not stable in acid or neutral solution. For example, the p-aminoethylbicarbonate ester is formed only in basic solutions of pH greater than 11. Since the pH of ethynolamine solutions is less than 10, the formation of these esters can be neglected (Astarita et al., 1983). The hydroxyl functional group may be considered inert for the systems employed in this study. For primary and secondary amines, the carbamates reaction channel provides a faster reaction rate than is available with tertiary amines, which can only react through formation of bicarbonates (or the formation of a Lewis acid-base adduct). As a consequence, tertiary amines are not employed for CO,. This fact, on the other hand, makes tertiary amines attractive as a solvent for mixtures of H2S, COz, and/or COS impurities in a gas stream. The removal of H2S is much faster than the removal of CO,, enabling significant chemical selectivity toward HzS under normal absorber operating conditions. DIPA is believed to offer better chemical selectivity because its structure inhibits the rate of the reaction that leads to carbamate formation. Steric hindrance, or partial blockage of the reactive lone pair electron site of the N by the bulky isopropyl groups, slows the rate of carbamate synthesis (Sigmund et al., 1981). The CO, reactions with aqueous solutions of MEA and DEA have been investigated extensively under a variety of experimental conditions. Batt (1979) has shown that water is not essential for carbamate formation when MEA and DEA are employed. Hikita et al. (1977) studied the kinetics of the reaction of COz(aq) with MEA and DEA solutions by using a rapid mixing thermal method which circumvented the problems of mass transfer a t a gas interface. The concentration of MEA and DEA was varied from 0.177 to 0.719 M over the temperature range 5.6-40.3
-
+
-
OC, and the C02/amine loadings were always below 0.5. In this concentration regime the rate of the reaction was found to be first order with respect to both COz and MEA. For DEA, the rate was found to be second order with respect to DEA. Danckwerts and Sharma (1976) presented the following rate expression for the reaction of C02with primary and secondary amines Rco2 = ~ A M ( A M ) ( C O ~ ) (9) where k A M is the second-order rate constant, (AM) is the concentration of amine, and (CO,) is the concentration of COz. Danckwerts (1970) reported the rate constants at 25 "C for DIPA to be 400, that for MEA to be 7600, and that for DEA to be 1500 M-' s-l. The only tertiary amine for which rate data are readily available is TEA. Hikita et al. (1977) proposed an ester forming reaction between COz and TEA as (HOCZHJ,N + COZ (HOC2HJNC2HdOCOOH (10)
-
They found the second-order rate constant for this reaction to be 50 M-l s-l from their kinetic studies, in which both the concentrations of CO, and TEA were varied in the temperature range 10-40 "C. However, in view of the known instability of carbonic acid esters at pH levels below 10, reaction 10 probably does not represent the reaction which occurs. Sada et al. (1976) determined the stoichiometry of the reaction between COz and TEA and observed that 1mol of CO, reacted with 2 mol of TEA. They proposed the reaction scheme (HOC2Hd)SN + HzO (HOCzH4)3NH++ OH(11) -+
CO,
+ 20H-
-
C032-+ H,O
(12)
This adds up to give the observed stoichiometric equation COS + 2TEA
+ H,O
+
(TEA)H+ C03,-
(13)
Sada et al. (1976) performed kinetic experiments using a wetted wall gas chromatographic column and report a second-order rate constant to be 16.8 M-l s-l at 25 "C. However, reaction 12 is known to be rapid, and the rate data do not allow reaction 11 to be the rate-determining step, so this scheme cannot be correct. The rate constants observed by both Sada et al. and Hikita et al. are very small. An alternate mechanism is that TEA undergoes the simple carbonate mechanism provided by reactions 4-7, which can also account for the observed stoichiometry. However, on the basis of the research reported in this paper, the formation of a Lewis acid-base complex followed by a base-catalyzed hydrolysis is also a possibility. In any case, the reaction between TEA and C02 deserves further investigation. In this research we extend the study of Batt (1979) and report the NMR spectra of the products of the interaction of H2S and CO, with anhydrous DGA and DIPA and interpret the observations of the effects of interacting these acid gases with MDEA and DMEA.
Experimental Section The experimental apparatus is shown in Figure 1. Rahman (1984) has complete details of the experimental apparatus and techniques. The reaction vessel was a Claisen distillation flask. The experimental volume of the reaction system with thermometer, gas buret, and supply lines inserted as in Figure 1 was 552.5 mL. Amine was added to the system through the capillary nozzle outlet of the high precision buret. The uncertainty in the amine volume was 0.025 mL. The temperature of the gas was measured at a point just above the gas-liquid interface. A mercury thermometer,
Ind. Eng. Chem. Res. Vol. 26, No. 1, 1987 29
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Figure 3. 13CNMR spectrum of pure DGA in DCC13 (20% sample).
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Figure 4. 'H NMR spectrum of DGA-H2S reaction products in DCC13 (20% sample). I
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Figure 2. 'H NMR spectrum of pure DGA in DCC13 (20% sample).
range -4 to 220 OF, was used to measure the gas temperature as the reaction progressed. When the reaction appeared to stop, as determined by a constant C02pressure, an aliquot of the reaction mixture was pipetted from the flask and diluted approximately by a factor of 5 for NMR analysis using either D20 or DCC1,. The piping consisted of 'I8-in. stainless steel, except for the line to the vacuum pump which was ll4-in. rubber tubing with lI4-in. walls. MEA and MDEA were obtained from Alfa Products and were stated to be 96 % and 97 % pure, respectively. Aldrich Chemical Company provided DMEA of 99% stated purity. Reagent grade DEA, assayed to be 99.8% pure, was obtained from Baker Chemical Company. ICN Pharmaceuticals supplied technical grade DIPA and DGA, both about 95% pure. Carbon dioxide and hydrogen sulfide gases were supplied by Linde and were 99.5% pure. Wilmad Glass Company supplied the reagents and sample tubes for NMR spectroscopic analysis. Deuterium oxide was 99.98% isotopically pure. Calibration of the NMR scale was performed using NMR grade tetramethylsilane, 99.9% pure, and perdeuteriotetramethylsilane, 99.9% pure.
Reaction between Hydrogen Sulfide and Diglycolamine The 'H and 13CNMR spectra of pure DGA in DCC13 are presented in Figures 2 and 3, respectively. The structural formula is
0 7 6 H2N8H2CH20CH2CH20H 1
The proton absorption peak labeled 1 in Figure 2 is attributed to proton impurity in DCC13. The set of absorption peaks labeled 2 are assigned to the @ and y pro-
tons, respectively. Absorption peak 4 is the response of the 6 protons, and the NH2 and OH proton resonances are assigned to absorption peak 3. In the 13Cspectrum, Figure 3, the similarity of the environments of the p and y carbon nuclei results in two absorption peaks at almost the same position, peaks 5 and 6. Peaks 7 and 8 are the 6 and cr carbon nuclei, respectively. When H2S was reacted with DGA and an aliquot subjected to NMR analysis at concentrations in DCCl,, NMR spectra almost identical with those given in Figures 2 and 3 were observed. In the case of the proton spectrum, the peaks appeared to be shifted down by 0.2 ppm. The absorption peaks labeled 9-12 correlate with those labeled 2-4 on Figure 2. Proton absorption by H2S, expected between 2 and 3 ppm, was either masked by the DGA absorptions or was too small for identification. The equilibrium reaction for the interaction of H2Swith anhydrous DGA is postulated to lead to a Lewis acid-base complex: H2S + H,N(CH,)20(CH2)20H HSH:NH,(CH,),O(CH2)20H (14) Since the concentration of H2S was not high enough to yield visible H2S absorption in the proton absorption spectrum, the failure to detect absorption in the proton absorption spectrum, the failure to detect the Lewis acid-base complex in either the IH or the 13C spectrum of the reaction products is not surprising. Since aqueous DGA is effective at H2S removal commercially, and the pressure measurements in this research indicated a chemical reaction has occurred, eq 14 must be readily reversed in an anhydrous DGA. When the reaction system was let down to atmospheric pressure for NMR analysis, the anhydrous amine was observed to yield gas which escaped. 1--*
Reaction between Carbon Dioxide and Dig1ycolamine When C02 was reacted with anhydrous DGA, the reaction product was highly viscous and the 'H NMR
30 Ind. Eng. Chem. Res. Vol. 26, No. 1, 1987
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Figure 7. 'H NMR spectrum of DIPA-H2S reaction products in DCC1, (10% sample).
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Figure 5. 13C NMR spectrum of DGA-C02 reaction products in DCC13 (20% sample).
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Figure 6. 'H NMR spectrum of pure DIPA in DCC1, (10% sample).
spectrum of the reaction products was of very poor resolution. However, a high-resolution 13Cspectrum was obtained and is presented in Figure 5 . The 13C resonance of the carbamate moiety was observed at 165 ppm. The assignments for peaks 14, 17, and 19 are the same as for peaks 6, 7, and 8, in pure DGA, Figure 3. Peaks 15, 16, and 18 as well as peak 13 must be attributed to the carbamate ion-pair product. Peaks 15 and 18 are tentatively assigned to the @ and a carbon nuclei of the carbamate product. The ratio of peak 18 to peak 19 is consistent with experimental conversion estimated from the COPuptake in the reaction flask. The structure of the product salt can be written as
HO(CH2)2O(CHz)2NH3+-OZCNH~(CH2)2O(CH2)zOH 2 The viscosity of the solution is consistent with an appreciable concentration of the ion-pair reaction product 2 in solution.
Reaction between Hydrogen Sulfide and Di-2-pro panolamine
The IH NMR spectrum of di-2-propanolamine (DIPA) in DCC13 is presented in Figure 6. The symmetric structure of this compound, i.e., HN[CHzCH(OH)CH3]2 3
permits reference to proton resonance in terms of the number of carbon atoms distant from the N atom, a referring to CH2protons and assigned to peak 22, p referring to CH of CH(0H) and assigned to peak 21, and X referring to the CH3protons and assigned to peak 23. The N-H and two 0-H absorptions are assigned to structureless peak 20. When H,S is reacted with DIPA, reaction 15, analogous to reaction 14, is anticipated. Since the freezing point [CH3CH(OH)CH2]2HN: + H2S [CH3CH(OH)CHZ]HN:HSH (15)
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Ind. Eng. Chem. Res. Vol. 26, No. 1, 1987 31
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Figure 9. I3C NMR spectrum of DIPA-C02 reaction products in DCC13 (10% sample).
H2Sfor the lone pair electrons on the N atom. The failure to observe any new peaks in the 13C NMR spectrum of the reaction products suggests that the concentration of the Lewis acid-base complex is too low to be observed or that the resonances are not resolvable from those of the pure amine.
Reaction between Carbon Dioxide and Di-2-propanolamine The lH NMR spectrum of the reaction products from the interaction of C02 and DIPA was indistinguishable from that presented for pure DIPA (Figure 6) and, hence, is not presented. Because the proton absorptions are largely unresolved in pure DIPA, gross changes in the proton spectrum of the products would be necessary for detection. However, the carbamate of DIPA was detected in the 13C spectrum of the reaction products, Figure 9. The salt-forming reaction is C02 + 2(CH3CHOHCH2)2NH
caping from the liquid amines. In view of the low concentrations of Lewis acid-base adducts for HzS and DIPA, which was solid, this result is not altogether surprising. Since carbamate formation is not possible for tertiary amines and water is not present to allow bicarbonate formation, the failure to observe any difference between the NMR spectra of these tertiary amines before and after C02simply indicates that the Lewis acid-base complex is too low in concentration for detection after venting to the atmosphere. There were small shifts in the locations of the proton and 13Cabsorptions for these tertiary amines, but in view of the fact that such shifts can be affected by concentration changes, the authors are loathe to attribute these to complex formation. These observations suggest that water plays a very important role in stabilizing the Lewis acid-base complex for both H2S and C 0 2 .
Conclusions The formation of carbamate from the interaction of COP and DGA and DIPA extends the observations of Batt (1979) to these amines and proves that water is also not required for carbamate formation in these cases. Evidence obtained from the NMR spectrum of the reaction products between H2S and DIPA suggests a Lewis acid-base complex is formed in the absence of water. However, no evidence for such a complex was found in the NMR spectrum of the products obtained from HzS and DGA. In the case of the tertiary amines, no evidence of chemical reaction was obtained from the NMR spectrum for either H2S or COP Pressure measurements suggest that Lewis acid-base complexes are formed with these liquid amines, but the complex decomposed when samples were taken from the reactor.
Registry No. 1,929-06-6; 2,105089-53-0; H2S,7783-06-4; COP, 124-38-9; diisopropanolamine, 110-97-4; methyldiethanolamine, [(CH3CHOHCH2)2NC02-][+H2N(CH2CHOHCHJ2] 105-59-9; dimethylethanolamine, 108-01-0; DIPA carbamate salt, 105102-40-7. (16) +
The /3 carbon nuclei are represented by the absorptions labeled 36 and 37, the CY carbon nuclei by absorptions labeled 39 and 40, and the X carbon nuclei by absorptions labeled 41 and 42, respectively, for both the protonated and the unreacted DIPA. The CY carbon nuclei and the /3 carbon nuclei of the carbamate are assigned to the absorptions labeled 38 and 35, respectively. The peak labeled 34 is tentatively assigned to absorption by the carboxyl carbon nucleus of the carbamate but is further upfield than anticipated. (The triplet at 77 ppm in this spectrum arises from HCC1, impurity in the solvent and appears to various degrees in all the spectra utilizing deuteriochloroform as the solvent.)
Reactions Involving Tertiary Amines Evidence for reaction products from the interaction of both H2Sand C02with anhydrous methyldiethanolamine (MDEA) and anhydrous dimethylethanolamine (DMEA) could not be found by comparison of the 'H and I3C NMR spectra for reacted and unreacted samples. Significant pressure drops were observed when these amines were introduced into the reaction flask containing either C02 or H2S. However, when the pressure was released for sampling for NMR analysis, bubbles were observed es-
Literature Cited Astarita, G.; Bisio, A.; Savage, D. W. Gas Treating With Chemical Soluents; Wiley-Interscience: New York, 1983. Astarita, G.; Savage, D. W.; Longo, J. M. Chem. Eng. Sci. 1982,37, 1953.
Batt, W. T. M.S. Thesis, Oklahoma State University, Stillwater, 1979.
Danckwerts, P. V. Gas-Liquid Reactions; McGraw-Hill: New York, 1970.
Danckwerts, P. V.; Sharma, M. M. Chem. Eng. 1966,202CE, 244. Hikita, H.; Asai, S.; Katsu, Y.; Ikuno, S. AIChE J. 1979,25, 793. Hikita, H.; Isikawa, H.; Honda, M. Chem. Eng. J . 1977,14, 27. Ladda, S. S.; Danckwerts, P. V. Chem. Eng. Sci. 1981,36,479. Maddox, R. N. Gas and Liquid Sweetening, 3rd ed.; Campbell Petroleum Series; Campbell Petroleum: Norman, OK, 1982. Nguyen, Y. N. Ph.D. Thesis, University of Rochester, Rochester, NY, 1979. Rahman, M. A. Ph.D. Thesis, Oklahoma State University, Stillwater, 1984. Sada, E.; Kumazawa, H.; Butt, M. A. Can. J . Chem. Eng. 1976,54, 421. Sharma, M.M.; Danckwerts, P. V. Chem. Eng. Sci. 1963,18, 729. Sigmund, P. W.; Butwell, K. F.; Wussler, A. J. Hydrocarbon Process. 1981,60,119.
Received for review August 9, 1983 Accepted April 1, 1986
