Xnd. Eng. Chem. Fundam. 1983,22, 239-249 Israelachvili, J. N.; Mitdrell, D. J.; Nlnham, B. W. J . Chem. Soc., Faraday Trans.2 1978, 72, 1525-1588. Kale, K. M.; Cussler, E. L.; Evans, D. F. J . Phys. Chem. 1980, 84.593-598. Kamrath, R. F. M.S. Thesls, Pwdue University, 1981. Kamrath. R. F.; Frames, E. I. submmed for publication In J . Colbid Interface Scl. (1982). Kamrath, R. F.; Fransea, E. I. I n Proceedingsof International Symposium of Surfactants In Solution, Lund, Sweden, June 27July 2. 1982, MHtai, K. L., Ed.; to be published by Plenum Press, 1983. Lange, H.;Beck, K. H. KoHOMZ.-Z. Polym. 1973, 251, 424-431. Morol. Y.; NlshikMo, N.; Matuura, R. J . ColbH Interface Scl. 1974. 46, 111-1 17. Morol, Y.; NishikMo, N.; Matuura. R. J . colldd Interface Sci. 19751, 50, 344-351. Moroi, Y.; NishkMo, N.; Salto. M.; Matuura, R. J . CoIbH Interface Sci. 1975b, 52. 358-383. Mukerjee, P. A&. ColbH Interface Sci. 1987, 1, 241-275. Mukerjee, P. I n “Physical Chemistry: Enriching Topics from Colloid and Surface Science”; IUPAC, van Olphen, H.;Mysels, K. J., Ed.; Theorex: La Wla, CA. 1975; pp 135-153. Mukerjee, P.; Mysels, K. J. I n “Colloidal Dlsperslons and Micellar Behavior”; MHtai, K. L.. Ed.; ACS Symposium Serbs No. 9, American Chemical Society. Washington, DC, 1975; pp 239-252.
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Mukerjee, P.; Yang, A. Y. S. J . Phys. Chem. 1978, 80, 1388-1390. Murray, R. C.; Hartley, 0.S. Trans. Farahy Soc. 1935, 31, 183-189. Mysels, K. J. J . Couold Interface Sci. 1978, 66, 331-334. Mysels, K. J.; Otter, R. J. J . Colbid Sci. 1981, 16, 482-473, 474-480. Perron, G.; Delisl, R.; Davldson, I.; Qenereux, S.; Desnoyers, J. E. J . Col/oH Interface Sci. 198i, 79, 432-442. Prausnk, J. M. “Molecular Thermodynamics of FiuibPhase Equlllbrla”;Prentice-Hail: Englewood Cllffs. NJ, 1989; pp 193, 289. Prigogine. I.; Defay, R. “ChemicalThermodynamics”;Longsmans (;reen and Co.: London, 1954. Rubingh, D. N. In ”Solution Chemlstry of Surfactants”; Vol. I , M h i , K. L., Ed.; Plenum Press: New York, 1979 pp 337-354. Shinoda. K. J . fhys. Chem. 1954, 58, 541-544. Shinoda, K.; Nomura, T. J . Phys. Chem. 1880, 84, 385-369. Stainsby. G.; Alexander, A. E. Trans. Faraday Soc. 1950, 46, 587-597. Stlgter, D. J . Colloid Interface Sci. 1974, 50. 473-482. Tanford, C. “the Hydrophoblc Effect”, 2nd ed.; Wiiey: New York, 1980.
Received for reuiew October 5, 1981 Revised manuscript receiued November 15, 1982 Accepted January 28, 1983
Sterically Hindered Amines for CO, Removal from Gases Guldo Sartorl’ and David W. Savage Corporate Research, Exxon Research and Engineering Company, Linden, New Jersey 07036
Steric hindrance and basicity are shown to control C0,-amine reactions. In aqueous amino alcohols, steric hlndrance is the dominant factor giving high thermodynamic capacity and fast absorption rates at high C02 loadings. Introducing steric hindrance by a bulky substituent adjacent to the amino group lowers the stability of the carbamate formed by C02-amine reactlon. Reduced carbamate stabiiity allows thermodynamic COP loadings to exceed those attainable with conventional, stable-carbamateamines. Lowering carbamate stability also leads to high free-amine concentration in solution; therefore fast amIne-CO, reaction rates are obtained despite some reduction of the rate constant owing to steric interference. Hindered amines show capacity and absorption rate advantages over conventional amines for COPremoval from gases by absorption in aqueous amine solutions and amine-promoted hot potassium carbonate. Cyclic capacity broadening of 20-40% and absorption rate increases up to 100% or more are possible with certain hindered amines.
Introduction Most commercial processes for the bulk removal of COz from gaseous streams involve the use of amines, usually amino alcohols (Kohland Reisenfeld, 1979). Amines are used either as an aqueous solution or as promoters for aqueous potassium carbonate solution. The choice of type of process depends primarily on the partial pressure of C02 in the feed gas and on the level of C02 desired in the treated gas (Tennyson and Schaaf, 1978). Amine-based processes are generally used at C02 partial pressures in the feed gas up to 1W200 psia. At higher pressures physical absorption in polar organic solvents may be preferred. The common aqueous amino alcohol processes (monoethanolamine (MEA), diethanolamine (DEA), diisopropanolamine (DIPA) and &@’-hydroxyaminoethylether (DGA))d show limited thdrmodynamic capacity to absorb C02. This can be seen in Figure 1, which presents vapor-liquid equilibrium data measured at 40 “C (Isaacs et al., 1974;Kent and Eisenberg, 1976; Lee et al., 1976). The reluctance of the common amino alcohols to load up with C02 much beyond 0.5 mol of C02/mol of amine can be attributed to the rather stable carbamates formed with amines in which the amino nitrogen is attached to a primary alkyl group C02 + 2RR’NH + RR’NCOO- + RR’N’H, (1) 0196-4313/83/1022-0239$01.50/0
When carbamate formation is the only reaction (eq l),the maximum loading is limited by stoichiometry to 0.5 mol/mol of amine. A certain amount of carbamate hydrolysis occurs with all amines so that even with MEA and DEA the loading may exceed 0.5, particularly at high pressures. Hydrolysis generates free amine which can react with additional C02, thus allowing the loading to exceed 0.5. A different limiting behavior is shown by tertiary amines which are unable to form carbamates. Aqueous tertiary amine reaction with C02 leads to formation of the bicarbonate ion C02 + R’R’’R”N
3
HCOB-
7
+ R’R’’R’’N+H
(2)
and stoichiometric absorption of COz can now reach 1.0 mol of C02/mol of amine. This is well supported by available experimental data, as shown in Figure 1 for methyldiethanolamine, MDEA (Jou et al., 1981). While the high C02pickup possible with tertiary amines is very attractive in practice, the low rates of COz absorption in tertiary amine solutions may limit their use. This paper describes the discovery of a new class of amines-sterically hindered amines-which approach the high thermodynamic capacity of 1 mol of C02/mol of amine combined with absorption rates comparable to those @ 1983 American Chemical Society
240
Ind. Eng. Chem. Fundam., Vol. 22, No. 2, 1983 Table I. Thermodynamic Constants for Common Amino Alcohols (Jensen et al., 1954; Chan and Danckwerts, 1981) (25 PK,,C) (20 KS C) monoethanolamine, MEA diethanolamine. DEA diisopropanolamine, DIPA triethanolamine, TEA methyldiethanolamine, MDEA
,a. = concentration of dissolved C02 in the liquid bulk @,/IT), and pi and p e = partial pressure of COz at the interphase and in the bulk, respectively. Unpromoted carbonate at 90 "C shows a mass-transfer enhancement factor in the range 10-30 (relative to physical absorption with a nonreactive solution having the same physical properties as K2C03solution). C02mass transfer in hot carbonate solution without promoters has been studied theoretically and experimentally (Savage et al., 1980). Addition of the conventional amine promoter, DEA, leads to significant rate promotion relative to unpromoted carbonate (about 300%) at low COz loadings. However, the promotional effect becomes smaller as the C02loading increases. Rate promotion results from the relatively high concentration of free amine in the carbonate solution compared to the concentration of hydroxyl ion. With the hindered amine promoter the rate promotional effect is large even relative to the DEA-promoted carbonate. This is due to both the high concentration of free amine (caused by low carbamate stability) and the large specific rate constant of highly basic amines. In fact, our analysis of amine promoters leads to the conclusion that COPabsorption rate with the most efficient amines at high temperature occurs in the regime where the chemical kinetics are essentially instantaneous compared to the time scale for diffusion in the solution. Thus, hindered amines offer the highest possible rate enhancement, and further rate enhancement would only be possible if the amine gave larger capacity enhancement than the present promoters give. Technological Impact COz scrubbing from gases by absorption with reaction in aqueous amino alcohol solutions employs absorption at high COzpartial pressure and low temperature (40-80 "C) followed by regeneration at reduced pressure and high temperature (120 "C) with steam stripping. In cycles of this type the hindered amine, AMP, offers a larger thermodynamic cyclic capacity than MEA. The absorption rate data for the hindered amines show that the full thermodynamic capacity can be utilized with fast rates. For C02 removal with amine-promoted carbonate solutions the absorption and regeneration steps operate at approximately the same temperature (about 110 "C) with regeneration by partial pressure reduction. In cycles of this type the hindered amine-promoted system offers broadened thermodynamic cyclic capacity. Furthermore, hindered amine promoters offer superior absorption rate enhancement when compared to conventional, linear amine promoters. In both types of processes, the use of sterically hindered amines leads to reduction in scrubbing solution rate and reduced energy requirement for solution regeneration. Conclusions The chemistry of COz reactions with amines has been shown to be controlled by steric hindrance and basicity. In aqueous amino alcohol solutions, steric hindrance is the dominant factor for achieving both high thermodynamic capacity and fast absorption rates at high C02 loadings. Steric hindrance is introduced by placing a bulky substituent group adjacent to the amino nitrogen site. The bulky substituent lowers carbamate stability. This favors carbamate reversion to bicarbonate ion and free amine leading to thermodynamic loadings approaching one mole of C02 per mole of amine. Hindered amines show capacity and absorption rate advantages over conventional, linear amines for COz re-
moval from gases by absorption in aqueous amines, aqueous polar organic solvent solutions of amines, and in amine-promoted potassium carbonate solutions. Acknowledgment We especially wish to acknowledge the contributions of Dr. M. T. Melchior in 13C NMR spectroscopy and Professor G. Astarita in thermodynamic analysis. We also wish to acknowledge the contributions of many of our colleagues who contributed directly and indirectly to the hindered amine science program: J. N. Begasse, M. Berger, A. L. Bisio, G. R. Chludzinski, B. Eisenberg, E. R. Ernst, P. M. Findeis, F. R. Gamble, H. C. Hayworth, J. N. Iyengar, R. R. Johnson, C. J. Kim, F. Leder, J. P. Longwell, G. E. Milliman, A. L. Pozzi, R. A. Reitz, and G. R. Say. Appendix In an aqueous amine solution of amine molarity m containing some chemically combined COz, it is possible to write m = [R2NH] + [R2NCOO-] + [RzNH2+] ( A l ) my = [RzNCOO-] + [HC03-] + [COS2-] (A2) where y represents the chemical loading in the solution expressed as moles of chemically combined C02 per mole of amine. The requirement of electrical neutrality implies that [R2NH2+]= [R,NCOO-]
+ [HCOc] + 2[C032-]
(A3)
Let K,, represent the second dissociation constant of carbonic acid. Then
Kc2at room temperature is 10-10.26.It follows that the ratio Kc2/K, is always a very small number (see Table I for values of pK,). This in turn implies that, with the exception of solutions of very basic amines at very low values of y , the COl- concentration can in general be regarded as negligible. If [COS2-]is taken as zero in eq A3, the system of eq 4, (text), Al, A2, and A3 can be solved to yield [RZNCOO-] S - [S2- 4y (1 - y)]'/' = fb) (A51 m 2
where
S = 1 + (l/Kcm); (1 5 S Im)
(-49) The function f(u) is symmetrical with respect to y = 0.5, i.e. fb)= f(1- Y) (A101 and its derivative is positive at y < 0.5 and negative at y > 0.5. Let K,, be the first dissociation constant for carbonic acid
and let H be the Henry's law constant for C02 physical
Ind, Eng. Chem. Fundam. 1983, 22, 249-258
solubility in the aqueous solution p* = H[C02]
(AW
The equilibrium vapor pressure of C02can be calculated from eq 3,4, A5-A8, and All-A12 as
Nomenclature AMP = 2-amino-2-methyl-1-propanol DEA = diethanolamine DIPA = diisopropanolamine MDEA = N-methyldiethanolamine MEA = monoethanolamine PE = 2-piperidine ethanol ai = concentration of dissolved COz at the interphase (pi/H) a. = concentration of dissolved COz in the liquid bulk (p,/H) fb)= (S- [SZ - 4y (1 - y)]'/2)/2 H = Henry's law constant for COP Ka = ([RZ"I[H+IJ/([R~NH~+I) k , = rate constant for reaction of C 0 2 with amine K c = [RzNC00-]/{ [RZNH] [HCOS-]) Kc, = W+I[HCO,-II/ [COP1 Kcz = W+l[co32-ll/ W O 3 - I k L = N / ( q - ao) k , = rate constant for zwitterion formation from amine and COZ m = total molar concentration of amine N = COz absorption flux per unit area p* = equilibrium vapor pressure of COz p e , pi = partial pressure of COz in the bulk and at the interphase, respectively s = 1 + (l/Kcrn)
249
y = moles of COP absorbed per mole of amine Registry No. COz, 124-38-9;KzC03, 584-08-7;HOCHzCHzNHCOZH, 6998-39-6;( H O C H ~ C H ~ ) ~ N C O 84836-10-2; ~H, HOCHZC(CH3)2NHC02H,84836-11-3; 2-amino-2-methy1-1-propano1, 124-68-5;2-piperidine ethanol, 1484-84-0. Literature Cited Astarita, 0. "Mass Transfer with Chemical Reaction"; Elsevier, 1967. Capiow. M. J. Am. Chem. SOC. 1868, 90, 6795. Chan, H. M.; Danckwerts. P. V. Chem. Eng. Sci. 1881, 36, 119. Danckwerts, P. V.; McNeil, K. M. Trans. Inst. Chem. Eng. 1967, 45,T32. Frahn, J. L.; Mliis, J. A. Aust. J. Chem. 1864, 17, 256. Isaacs, E. E.; Otto, F. D.; Mather, A. E. Can. J. Chem. Eng. 1874, 52, 125. Jensen, M. B.; Jorgensen, E.; Faurholt, C. Acta Chem. Scand. 1954, 8 , 1137. Jensen, M. B. Acta Chem. Scand. 1857, 1 1 , 499. Jou, F. Y.; Lal, D.; Otto, F. D.; Mather, A. E. "The Solubllity of H2S and COP in Aqueous Methyldiethanolamine Solutions", Paper 20e, presented at AIChE Meeting, Houston, Apr 5-9, 1981. Jones, J. H.; Fronlng, H. R.; Clayton, E. E., Jr. J. Chem. Eng. Data 1858, 4 , 85. Kent, R. L.; Eisenberg, 8. "Proceedings, Gas Conditioning Conference"; University of Okalahoma, 1975; Paper E. Kohl, F.; Reisenfeld, F. C. "Gas Purification", 3rd ed., Gulf Publishing: Houston, 1979. Lee, J. I.; Otto, F. D.; Mather, A. E. Can. J. Chem. Eng. 1878, 54,214. Melchior, M. T., Exxon Research and Engineering Company, Linden, NJ, private communication, 1977a. Melchior, M. T., private communlcatlon, 1975. Melchior, M. T., private communication, 1977b. Savage, D. W.; Astarita, G.; Joshi, S. Chem. Eng. Sci. 1880, 358 1513. Sharma, M. M. Trans. Faraday SOC. 1865, 61, 661. Tennyson, R. N.; Schaaf, R. P. Oil Gas J . 1877, 75(2), 78. Tosh, J. S.;Field, J. H.; Benson, H. E.; Haynes, W. P. "Equilibrium Study of the System Potassium Carbonate, Potassium Bicarbonate, Carbon Dioxide, and Water"; U S . Bureau of Mines Report of Investigations, No. 5484. 1959.
Received for review August 31, 1981 Revised manuscript received January 3, 1983 Accepted January 21, 1983
Thermodynamics of Highly Solvated Liquid Metal Solutions Montgomery M. Alger and Charles A. Eckert' Department of Chemical Engineering, University of Illinois, Urbana, Illinois 6 180 1
Chemical theory has been found to be well-suited for representing abrupt changes that are often observed in the activity coefficients and partial molar enthalpies of the components in systems that are known to form compounds in the solid phase. A chemical model is presented in which the Gibbs energy and the enthalpy are coupled through a simple analytical expression. The parameters that are obtained from the fitting of experimental data can be interpreted physically because of the rigorous thermodynamic foundation of the model. The model provides a rational method for extending limited thermodynamic data for a given system to a wider range of temperatures and compositions.
Introduction Many liquid metal solutions that form compounds in the solid phase do not conform to the regular solution theory requirement of zero excess entropy of mixing. Experimental entropies of mixing (Hultgren et al., 1973) often indicate the presence of unusual amounts of ordering in the liquid phase at compositions where compounds are known to exist in the solid phase. Jordan (1979) and Predel(l979) have presented evidence that suggests that compounds present in the solid phase also exist in the liquid phase when the alloy melts. Any attempt to model deviations from ideal mixing behavior in these systems should account for compound formation in the liquid phase. 0196-43 13/83/1022-0249$01.50/0
Eckert et al. (1981) have recently presented a general treatment of chemical theory as applied to strongly solvating (compound-forming) liquid metal systems. A comparison made between the Van Laar, Wilson, and three suffix Redlich-Kister equations and chemical theory showed that chemical theory gave a better fit of -y@ in the Mg-Bi system with the same number of parameters. In this paper a unified chemical theory model is presented. The model is capable of representing the Gibbs energy and the temperature dependence of the Gibbs energy in compound-forming systems with a minimum number of parameters. All the model parameters are based on well-defined thermodynamic functions. Because of the simple analytical form of the model, mathematical oper0 1983 American Chemical Society
