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Ind. Eng. Chem. Res. 1999, 38, 4032-4036
A Study on Equilibrium Solubility for Carbon Dioxide in Methyldiethanolamine-Piperazine-Water Solution Hua-Bing Liu,* Cheng-Fang Zhang, and Guo-Wen Xu Research Institute of Chemical Technology, East China University of Science and Technology, P.O. Box 274, 130 Meilong Road, Shanghai 200237, People’s Republic of China
Solubilities of CO2 in aqueous mixtures of methyldiethanolamine (MDEA) and piperazine (PZ) have been measured for temperatures and CO2 partial pressures ranging from 30 to 90 °C and 13.16 to 935.3 kPa, respectively. A modified Deshmukh-Mather thermodynamic model is used to correlate the experimental data with the average deviation of 9.8%, in which the effect of salts on Henry’s constant is taken into consideration. The results also indicate that the secondorder dissociation reaction for piperazine can be neglected. A simple model is also given with the average deviation of 11.6%. Introduction Activated methyldiethanolamine (MDEA) technology has been frequently applied for the removal of acidic gases such as CO2 and H2S from gas streams. Despite relatively high capital costs and expensive solvents, the advantages of flexibility, ease of operation, and lowenergy expenditure have made it attractive for new installations as well as revamps. Some investigations involving CO2 equilibrium solubility in the MDEA-based solution have been done. Chakravarty1 first suggested the use of aqueous blends of primary and tertiary amines for stripping CO2. He represented the solubility data of H2S and CO2 in aqueous mixtures of MDEA with monoethanolamine (MEA) or diethanolamine (DEA) with the model of Deshmukh-Mather,2 in which activity coefficients of solutes are calculated with the Guggenheim expression.3 Austgen et al.4 reported the CO2 solubility in 2.0 kmol/m3 MDEA + 2.0 kmol/m3 MEA or DEA aqueous solution at 40 and 80 °C for CO2 partial pressures below 315 kPa, and a thermodynamic model was developed by applying the electrolyte-NRTL (nonrandom two-liquid) activity coefficient equations. Shen and Li5 and Li and Shen6 measured the CO2 solubility in aqueous blends of MDEA and MEA at the total amine concentration of 30 wt %. Li and Shen7 expressed the chemical equilibrium constants involving amines as functions of temperature, amine concentration, and carbon dioxide solubility on the basis of the model of Kent-Eisenberg.8 Jou et al.9 determined the CO2 solubility in four aqueous mixtures of MDEA with MEA at temperatures from 25 to 120 °C over a range of CO2 partial pressures from 0.001 to 19930 kPa. Dawodu and Meisen10 provided the solubility data for CO2 in blends of 3.4 kmol/m3 MDEA + 0.8 kmol/m3 MEA or DEA and 2.1 kmol/m3 MDEA + 2.1 kmol/m3 MEA or DEA at temperatures from 70 to 180 °C and CO2 partial pressures from 100 to 3850 kPa. Xu et al.11 measured the solubility data for CO2 in 4.28 kmol/m3 MDEA with the piperazine concentrations ranging from 0 to 0.515 kmol/m3 and CO2 partial pressures from 3.83 to 76.77 kPa. * To whom correspondence should be addressed. Telephone: 0086-21-64252386. Fax: 0086-21-64250884. E-mail:
[email protected].
Piperazine (PZ) is an effective activator for an industrial CO2 removal process; however, the equilibrium data for CO2 in aqueous blends of MDEA with PZ is very scarce. In this study, the equilibrium solubility of CO2 in the MDEA-PZ-H2O solution was systematically determined and represented by a modified DeshmukhMather model and a simple model. Experimental Section The apparatuses used to determine the experimental data for this work have been reported in detail by Xu et al.11 and the experiment was carried out in two sets of apparatuses. For the CO2 partial pressures below 101.3 kPa, the experiment was conducted in an atmospheric pressure apparatus. In the case of CO2 partial pressures above atmospheric pressure, an elevated pressure apparatus consisting of a stainless steel autoclave was used to obtain equilibrium data. The experimental temperature was controlled to within (0.1 °C. Normally, it takes about 4 and 6 h for the atmospheric and elevated pressure system to reach equilibrium, respectively. The CO2 loading in the liquid phase was determined by volumetric analysis. The CO2 partial pressure was calculated by subtracting the partial pressure of water from the total pressure of the system. The equilibrium partial vapor pressure of water was investigated in aqueous 10-70 wt % MDEA solutions in our previous work and a part of the experimental data has been reported.11 It was found that the equilibrium water partial pressure obeys Raoult’s law well in an aqueous MDEA solution with the average deviation below 2.0%. In this work, Raoult’s law for water partial pressure is assumed to be also applicable to the MDEA-PZ-H2O system. The equilibrium partial pressures of MDEA and piperazine in the MDEA-PZ-H2O system have also been studied by Xu et al.11 and were found to be extremely low. Compared with the CO2 partial pressure, their partial vapor pressures were neglected with the error being no more than 1%. The Fifth Chemical Factory in Wujin of JiangSu province (China) supplied MDEA with +99% purity and piperazine is chemical grade with 98.8% purity. To correct the two experimental apparatuses and analytical procedures, the solubility of CO2 in 4.28 kmol/ m3 MDEA solution at 40 and 70 °C was measured. As
10.1021/ie990113v CCC: $18.00 © 1999 American Chemical Society Published on Web 09/15/1999
Ind. Eng. Chem. Res., Vol. 38, No. 10, 1999 4033
can be seen in Figure 1, our data are in good agreement with those reported in the literature.4,12 In this work, the CO2 solubility in the aqueous mixtures of MDEA with PZ was investigated the temperature range from 30 to 90 °C and the CO2 partial pressure from 13.16 to 935.3 kPa, the concentration of MDEA ranging from 1.53 to 4.77 kmol/m3, and that of PZ ranging from 0 to 1.55 kmol/m3. The experimental solubility data are listed in Table 1. Thermodynamic Model 1. Basic Relationship. The chemical reaction equilibria and gas-liquid-phase equilibria for the CO2MDEA-PZ-H2O system are coexistent. The main reactions in the system are as follows:
PZH+ S PZ + H+ +
Figure 1. Comparison of solubility of CO2 in 4.28 kmol/m3 MDEA solution at 40 and 70 °C between this work and the literature. 40 °C: ], Jou; 4, Austgen; b, this work. 70 °C: O, Jou; 9, this work.
+
MDEAH S MDEA + H
H2O + CO2 S H+ + HCO3-
Here, the first term represents the electrostatic effects of the solvent on the solute species at infinite dilution; the remaining term takes into account short-range van der Waals forces and βij is the interaction parameter for two solute species. The Debye-Hu¨ckel proportionality factor, A, can be calculated by
H2O S H+ + OHHCO3- S H+ + CO32Thus, the corresponding thermodynamic equilibrium constant expressions can be written
K1 ) γPZmPZγH+mH+/γPZH+mPZH+
(1)
K2 ) γMDEAmMDEAγH+mH+/γMDEAH+mMDEAH+
(2)
K3 ) γH+mH+γHCO3-mHCO3-/γCO2mCO2aW
(3)
K4 ) γH+mH+γOH-mOH-/aW
(4)
K5 ) γH+mH+γCO32-mCO32-/γHCO3-mHCO3-
(5)
In addition, mass and charge balances governing the reacting species can be formed:
mPZ + mPZH+ ) c1/F
(6)
mMDEA + mMDEAH+ ) c2/F
(7)
mCO2 + mHCO3- + mCO32- ) (c1 + c2)RCO2/F
(8)
mH+ + mPZH+ + mMDEAH+ ) mOH- + mHCO3- + 2mCO32- (9) Gas-liquid equilibrium for CO2 may be described as
φCO2yCO2P )
/ γCO2mCO2HCO F 2
(10)
The fugacity coefficient of CO2 was calculated by using the Peng-Robinson equation of state.13 As in the work of Deshmukh and Mather,2 the activity of water is set to its mole fraction and the activity coefficients of the solute species were calculated with the extended Debye-Hu¨ckel expression.3
ln γi ) -
Azi2I0.5 1 + I0.5
∑j βijmj
+2
(11)
A)
(
2πNAdw 1000
)( ) 1/2
e2 DkT
3/2
(12)
In this work, the density of water, dw, was calculated by using an expression given by Cheng et al.14 The dielectric constant of water, D, was calculated from the equation proposed by Maryott and Smith.15 Al-Ghawas et al.16 measured the physical solubility of CO2 in aqueous 0-50 wt % MDEA solutions for temperatures between 15 and 50 °C by the N2O analogy method. In this work, Henry’s constant correlation given by Al-Ghawas et al. is extended to the MDEA-PZ-H2O system and converted from a MDEA mass fraction basis to a molarity basis. As carbon dioxide undergoes chemical reactions with mixed solutions, MDEAH+, PZH+, HCO3-, HCO3-, and CO32- are formed. The increase in ionic strength of solution results in a decrease in gas solubility. Joosten and Danckwerts17 suggested an extended van Krevelen model for the mixed electrolyte solutions. Thereby Henry’s constant of CO2 in the aqueous mixture of MDEA with PZ is produced as follows: / /HCO2) ) log(H CO 2
∑hiIi
(13)
Here, hi and Ii are the van Krevelen coefficient and ionic strength attributable to the electrolyte i. Within this experimental condition, the concentrations of MDEAH+ and HCO3- are much higher than those of other ions (see Figure 3) and have predominant influence. Thus, the effects of the ions, except MDEAH+ and HCO3-, can be neglected. Reliable van Krevelen coefficients for MDEAH+ and HCO3-, have been experimentally determined by Browning and Weiland.18 The equilibrium constants and Henry’s constant needed for calculations are listed in Table 2. 2. Results and Discussion. There are two dissociation reactions for piperazine in an aqueous solution; thus, piperazine in an aqueous solution exists not only
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Ind. Eng. Chem. Res., Vol. 38, No. 10, 1999
Table 1. CO2 Equilibrium Partial Pressures of Experimental and Calculated Values t (°C) 50
70
50
70
30
50
70
90
30
50
70
90
50
70
50
70
aCO2
pexp (kPa)
pmod1 (kPa)
0.17 kmol/m3 PZ + 1.53 kmol/m3 MDEA 0.468 21.18 19.78 0.589 44.00 42.49 0.696 89.44 82.44 0.851 271.9 238.2 0.980 669.4 702.5 0.387 35.43 34.04 0.492 71.29 69.07 0.641 148.8 171.9 0.811 418.8 486.3 0.876 688.8 746.0 0.35 kmol/m3 PZ + 1.35 kmol/m3 MDEA 0.500 17.78 18.89 0.609 41.14 38.32 0.707 89.43 72.23 0.936 509.4 415.4 0.955 586.9 493.9 0.349 17.60 16.82 0.427 32.07 31.63 0.558 71.26 77.95 0.759 243.8 279.4 0.790 296.3 342.8 0.35 kmol/m3 PZ + 3.15 kmol/m3 MDEA 0.477 16.73 19.41 0.642 54.24 62.42 0.714 97.47 108.0 0.812 247.5 253.6 0.842 407.5 342.4 0.377 23.95 26.23 0.455 42.97 45.60 0.573 90.11 99.49 0.665 200.1 185.2 0.750 422.6 342.5 0.287 33.86 38.09 0.324 48.08 52.85 0.470 178.0 157.6 0.591 368.0 348.7 0.691 573.0 662.7 0.147 19.88 19.21 0.191 37.54 38.05 0.247 82.54 75.83 0.407 312.5 311.6 0.490 482.5 553.6 0.7 kmol/m3 PZ + 2.8 kmol/m3 MDEA 0.506 15.60 18.63 0.647 48.43 52.96 0.733 97.45 107.1 0.793 202.5 184.9 0.880 460.0 449.0 0.410 18.98 23.16 0.540 49.60 58.53 0.617 90.17 99.81 0.694 195.1 174.0 0.768 380.1 308.1 0.274 19.24 19.31 0.366 52.64 45.41 0.410 73.14 64.87 0.541 177.8 169.9 0.766 935.3 812.5 0.198 17.37 17.84 0.221 25.37 25.16 0.254 37.19 38.94 0.401 157.2 170.4 0.540 412.2 483.6 0.53 kmol/m3 PZ + 4.77 kmol/m3 MDEA 0.318 42.51 35.30 0.415 91.15 70.36 0.649 326.2 318.17 0.706 508.7 497.9 0.760 753.7 813.7 0.193 35.83 34.03 0.252 75.61 64.70 0.396 203.1 207.5 0.529 460.6 484.9 0.592 713.1 710.7 3 1.55 kmol/m PZ + 3.75 kmol/m3 MDEA 0.349 14.34 15.40 0.454 46.49 36.88 0.525 91.15 66.48 0.650 278.3 216.0 0.746 678.3 677.1 0.247 13.16 15.15 0.323 37.35 34.35 0.544 224.7 224.8 0.635 479.7 479.2 0.665 667.2 632.8
pmod2 (kPa) 18.26 40.91 81.17 244.5 809.5 27.80 59.61 159.6 506.9 867.3 18.21 37.58 71.29 467.4 584.1 14.80 28.07 72.01 279.5 350.4
Figure 2. Comparison of the experimental partial pressures of CO2 and thermodynamic model values.
22.21 65.71 104.2 206.5 261.7 27.72 49.90 111.4 203.2 357.4 35.10 48.93 155.6 356.6 687.3 27.02 37.67 65.25 272.2 507.3 21.67 55.06 97.42 148.5 301.8 26.63 69.49 117.3 197.0 328.9 47.90 19.67 69.67 191.9 918.4 17.85 24.73 37.78 174.2 531.3 31.59 69.89 350.2 510.0 735.3 26.54 46.57 168.7 451.4 687.1 19.11 48.50 85.32 213.0 424.5 13.10 33.79 267.1 528.7 662.0
Figure 3. Solute species concentration profiles in a CO2-loaded 0.7 kmol/m3 PZ + 2.8 kmol/m3 aqueous MDEA solution at 50 °C. [, HCO3-, 2, MDEAH+; b, MDEA; 0, PZH+; 4, PZ; ×, CO32-; O, CO2.
as the molecule PZ but also as PZH+ and PZH22+. The dissociation constants for piperazine are available only at 23 °C from the Handbook of Chemistry and Physics.19
PZH22+ S PZH+ + H+ PZH+ S PZ + H+
pK1 ) 5.56 pK2 ) 9.83
The concentrations of the three forms depend strongly on the hydrogen ion concentration. However, the pH of a carbonated alkanolamine solution of commercial application is normally in the range of 8-10. Simple analysis indicates that the ratio of PZH+/PZH22+ is in the range from 275 to 2.7 × 104 at 23 °C. It means that the second-order dissociation reaction for piperazine in the solution is nearly negligible. The first-order dissociation constant of piperazine, together with 14 specific interaction parameters, is determined by fitting the model to the measured CO2 equilibrium data. With these values in Tables 3 and 4, the vapor-liquid equilibrium of CO2-MDEA-PZ-H2O
Ind. Eng. Chem. Res., Vol. 38, No. 10, 1999 4035 Table 2. Expressions Used for the Thermodynamic Model expression
ref
ln K1 ) -22.086 + 459.37/T log K2 ) -14.01 + 0.0184T ln K3 ) 235.482 - 12092.1/T - 36.7816 ln T ln K4 ) 140.932 - 13445.9/T - 22.4773 ln T ln K5 ) 220.067 - 12431.7/T - 35.4819 ln T ln HCO2 ) a1 + a2/T + a3/T2 a1 ) 2.01874 - 2.83179cMDEA/F + 4.11932(cMDEA/F)2 - 0.81256(cMDEA/F)3 a2 ) 3135.49 + 1846.22cMDEA/F 2612.63(cMDEA/F)2 + 508.592(cMDEA/F)3 a3 ) -813702 - 295623cMDEA/F + 414660(cMDEA/F)2 - 79674.6(cMDEA/F)3
this work Barth et al.20 Edwards et al.21 Edwards et al.21 Edwards et al.21 Al-Ghawas et al.16
Thus, only three main dissociation reactions are taken into consideration and are expressed as follows:
K1′ ) cPZcH+/cPZH+
(1)
K2′ ) cMDEAcH+/cMDEAH+
(2)
K3′ ) K3′′cH2O ) cHCO3-cH+/cCO2
(3)
The balance relations can be approximated as follows:
cPZ + cPZH+ ) c1
(4)
cMDEA + cMDEAH+ ) c2
(5)
cHCO3- ) (c1 + c2)RCO2
(6)
cPZH+ + cMDEAH+ ) cHCO3-
(7)
Table 3. Specific Interaction Parameters for CO2-MDEA-PZ-H2O parameter
kg/mol
β(HCO3--PZH+) β(HCO3--MDEAH+) β(HCO3--PZ) β(HCO3--MDEA) β(PZH+-CO32-) β(PZH+-CO2) β(PZH+-PZ)
-0.1039 0.0385 -0.2248 -0.0049 -0.0877 -0.0778 0.0853
parameter
kg/mol
β(PZH+-MDEA) 0.0939 β(MDEAH+-CO32-) -0.0372 β(MDEAH+-CO2) -0.0056 β(MDEAH+-PZ) -0.1051 β(MDEAH+-MDEA) 0.0482 β(CO32--PZ) 0.4041 β(CO32--MDEA) 0.0872
A gas-liquid equilibrium of CO2 is represented as follows:
PCO2 ) HcCO2
Table 4. Expressions Used for the Simple Model expression
(8)
ref
ln K3 ) -241.828 + 29.8253 × 1.48528 × 108/T2 + 0.332647 × 11 3 13 4 10 /T - 0.282393 × 10 /T ln H ) 20.2669 - 1.38306 × 104/T′ + 0.06913 × 108/T2 - 0.015589 × 1011/T3 + 0.012 × 1013/T4 104/T′
Kent and Eisenberg8
From eqs 1-8, a direct relationship between the CO2 loading and its partial pressure is established:
Kent and Eisenberg8
can be calculated. The model values of CO2 partial pressure are listed in Table 1 with the average deviation of 9.8%. Figure 2 shows that the calculated results are satisfactory. The chemical equilibrium behavior of the carbonated solution, which varies with the CO2 loading, is also investigated. Figure 3 illustrates the concentration profiles of solute species in 0.7 kmol/m3 PZ + 2.8 kmol/ m3 MDEA solution as a function of CO2 loading at 50 °C. Within the experimental range of CO2 partial pressures, the concentrations of both MDEAH+ and HCO3- rise rapidly and remarkably above those of other cations and anions, which is accompanied by a sharp drop in the concentration of free MDEA. A slow increase in the concentration of PZH+ and a slow decrease in that of free PZ can be observed. Although the absorbed CO2 exists mainly in the form of bicarbonate, it shows that the content of CO32- cannot be ignored, and that of CO2 also becomes important at high loadings. A representation of equilibrium concentration profiles of solute species in the system not only is helpful for understanding vapor-liquid equilibrium behavior but also makes our constant calculation convincible. A Simple Model To make calculation convenient and save computation time, a simple model is also proposed. The concentrations of H+ and OH- are rather low, so it is reasonable to neglect their effects on the mass- and charge-balance equations. Another assumption is that all forms of the absorbed carbon dioxide are regarded as bicarbonate since the contents of CO32- and CO2 are also very low.
PCO2 H
)
RCO2
(-B + xB2 - 4AC)
1 - RCO2
2K3′
(9)
where
A ) (c1 + c2)(1 - RCO2)
(10)
B ) c1K2′ + c2K1′ - (c1 + c2)(K1′ + K2′ )RCO2
(11)
C ) -(c1 + c2)K1′ K2′ RCO2
(12)
In this simple model, the nonideality of the system is lumped into equilibrium constants of the main reaction, which are expressed as a function of temperature, CO2 solubility, and absorbent concentration. Using the leastsquares fit to the equilibrium partial pressure data of CO2 in aqueous MDEA + PZ solutions, two equilibrium constants governing the protonation reactions of MDEA and PZ are determined as follows:
ln K1′ ) -21.132 + 3.912 × 102/T - 3.865 × 107/T3 - 4.709 × 10-2/(1 - RCO2) - 5.875 × 10-2/RCO2 - 6.698 × 10-3/RCO22 - 0.5048 ln c1 (13) ln K2′ ) -9.676 × 10-2 - 7.166 × 103/T + 1.097 × 108/T3 - 2.068 × 10-2/(1 - RCO2) - 4.754 × 10-1/RCO2 + 6.762 × 10-2/RCO22 - 6.038 ln c2 (14) The equilibrium constant and Henry’s law coefficient for CO2 in water are presented in Table 4. This semiempirical model can be solved directly without iterative calculation. The average deviation between the calcu-
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Ind. Eng. Chem. Res., Vol. 38, No. 10, 1999
Literature Cited
Figure 4. Comparison of the experimental partial pressures of CO2 and simple model values.
lated and experimental CO2 partial pressure data is 11.6% and good agreement can be found in Figure 4. Nomenclature aw ) activity of water A ) constant in the extended Debye-Hu¨ckel expression ci ) concentration of solute i, kmol/m3 c1 ) fresh PZ concentration, kmol/m3 c2 ) fresh MDEA concentration, kmol/m3 dw ) density of water, kg/m3 D ) dielectric constant of water e ) elementary charge h ) van Krevelen coefficient H ) Henry’s law constant of CO2 in pure water, kPa‚m3/ kmol HCO2 ) Henry’s law constant of CO2 in amine solution, atm‚ m3/kmol / ) Henry’s law constant of CO2 in the loaded soluH CO 2 tion, atm‚m3/kmol I ) ionic strength k ) Boltzmann constant K, K′, K′′ ) equilibrium constant m ) mol/kg of water NA ) Avogadro’s constant P ) total pressure, atm p ) CO2 partial pressure, kPa t ) temperature, °C T ) temperature, K y ) gas-phase molar fraction z ) ionic charges Greek Symbols RCO2 ) CO2 loading, mol of CO2/mol of (MDEA + PZ) βij ) specific interaction, kg/mol γ ) molar activity coefficient φ ) gas-phase fugacity coefficient F ) density of water in solution, kg/L Subscripts i, j ) species or component mod 1 ) thermodynamic model mod 2 ) simple model Amine Abbreviations MDEA ) methyldiethanolamine PZ ) piperazine
(1) Chakravarty, T. Solubility Calculation for Acid Gases in Amine Blends. Ph.D. Dissertation, Clarkson University, 1985. (2) Deshmukh, R. D.; Mather, A. E. A Mathematical Model for Equilibrium Solubility of Hydrogen Sulfide and Carbon Dioxide in Aqueous Alkanolamine Solutions. Chem. Eng. Sci. 1981, 36, 355. (3) Guggenheim, E. A. The Specific Thermodynamic Properties of Aqueous Solutions of Strong Electrolytes. Philos. Mag. 1935, 19, 588. (4) Austgen, D. M.; Rochelle, G. T.; Chen, C.-C. Model of VaporLiquid Equilibria for Aqueous Acid Gas-Alkanolamine Systems. 2. Representation of H2S and CO2 Solubility in Aqueous Mixtures of MDEA with MEA or DEA. Ind. Eng. Chem. Res. 1991, 30, 543. (5) Shen, K.-P.; Li, M.-H. Solubility of Carbon Dioxide in Aqueous Mixtures of Monoethanolamine with Methyldiethanolamine. J. Chem. Eng. Data 1992, 37, 96. (6) Li, M.-H.; Shen, K.-P. Densities and Solubilities of Solutions of Carbon Dioxide in Water + Monoethanolamine + N-Methyldiethanolamine. J. Chem. Eng. Data 1992, 37, 288. (7) Li, M.-H.; Shen, K.-P. Calculation of Equilibrium Solubility of Carbon Dioxide in Aqueous Mixtures of Monoethanolamine with Methyldiethanolamine. Fluid Phase Equilibr. 1993, 85, 129. (8) Kent, R. L.; Eisenberg, B. Better Data For Amine Treating. Hydrocarbon Process. 1976, 55, 87. (9) Jou, F.-Y.; Otto, F. D.; Mather, A. E. Vapor-Liquid Equilibrium of Carbon Dioxide in Aqueous Mixtures of Monoethanolamine and Methyldiethanolamine. Ind. Eng. Chem. Res. 1994, 33, 2002. (10) Dawodu, O. F.; Meisen, A. Solubility of Carbon Dioxide in Aqueous Mixtures of Alkanolamines. J. Chem. Eng. Data 1994, 39, 548. (11) Xu, G.-W.; Zhang, C.-F.; Qin, S.-J.; Guo, W.-H.; Liu, H.-B. Gas-Liquid Equilibrium in a CO2-MDEA-H2O System and the Effect of Piperazine on It. Ind. Eng. Chem. Res. 1998, 37, 1473. (12) Jou, F.-Y.; Mather, A. E.; Otto, F. D. Solubility of H2S and CO2 in Aqueous Methyldiethanolamine Solution. Ind. Eng. Chem. Process Des. Dev. 1982, 21, 539. (13) Peng, D.-Y.; Robinson, D. B. A New Two-Constant Equation of State. Ind. Eng. Chem. Fundam. 1976, 15, 59. (14) Cheng, S.; Meisen, A.; Chakma, A. Predict Amine Solution Properties Accurately. Hydrocarbon Process. 1996, 75, 81. (15) Maryott, A. A.; Smith, E. R. Table of Dielectric Constants of Pure Liquids; NBS Circular 514; U.S. Government Printing Office: Washington, DC, 1951. (16) Al-Ghawas, H. A.; Hagewiesche, D. P.; Ruiz-Ibanez, G.; Sandall, O. C. Physicochemical Properties Important for Carbon Dioxide Absorption in Aqueous Methyldiethanolamine. J. Chem. Eng. Data 1989, 34, 385. (17) Joosten, G. E. H.; Danckwerts, P. V. Solubility and Diffusivity of Nitrous Oxide in Equimolar Potassium CarbonatedPotassium Bicarbonate Solutions at 25 °C and 1 atm. J. Chem. Eng. Data 1972, 17, 452. (18) Browning, G. J.; Weiland, R. H. Physical Solubility of Carbon Dioxide in Aqueous Alkanolamines via Nitrous Oxide Analogy. J. Chem. Eng. Data 1994, 39, 817. (19) Lide, D. R., Ed. Handbook of Chemistry and Physics, 77th ed.; CRC Press: Boca Raton, FL, 1996. (20) Barth, D.; Tondre, C.; Delpuech, J. J. Kinetics and Mechanisms of the Reactions of Carbon Dioxide with Alkanolamines: A Discussion Concerning the Cases of MDEA and DEA. Chem. Eng. Sci. 1984, 39, 1753. (21) Edwards, T. J.; Maurer, G.; Newman, J.; Prausnitz, J. M. Vapor-Liquid Equilibria in Multicomponent Aqueous Solutions of Volatile Weak Electrolytes. AIChE J. 1978, 24, 966.
Received for review February 16, 1999 Revised manuscript received May 20, 1999 Accepted May 22, 1999 IE990113V
