2006
Ind. Eng. Chem. Res. 1994,33, 2006-2015
Correlation and Prediction of the Solubility of Carbon Dioxide in a Mixed Alkanolamine Solution Yi-Gui Lit and Alan E. Mather' Department of Chemical Engineering, University of Alberta, Edmonton, Alberta, Canada T6G 2G6
The Clegg-Pitzer equations, expressed on a mole fraction basis, comprise an extended DebyeHiickel term which is a function of composition and a Margules expansion carried out to the threesuffix level, with some simplificationsare used to correlate and predict the vapor-liquid equilibrium data for the CO2-mixed amine aqueous system. The interaction parameters determined from data for the binary and ternary (single amine) systems can be used to predict the quaternary mixed amine system without any additional parameters. T h e model is applied to the C02-MDEA-MEA-H20 system containing three neutral solvents (MDEA, MEA, and H2O) and four ionic species (MDEAH+, MEAH+, HC03-, and RNHCOO-) with four chemical equilibria simultaneously involved. Introduction Aqueous alkanolamine solutions are frequently used for the removal of acidic gases, such as CO2 and H2S from industrial and natural gases. Among them, monoethanolamine (MEA, H2NCH2CH20H) solutions have been used extensively due to their high reactivity, low cost of the solvent, ease of reclamation, and low solubility of hydrocarbons. However, because of the formation of rather stable carbamate, the loading capacity of MEA solutions is lower than that of MDEA solutions. In recent years, MDEA (methyldiethanolamine, CH3N(CH&HzOH),) solution has been used as an alternative to MEA in certain gas-treating applications. MDEA has a higher loading capacity and low heat of reaction with the acid gases, which leads to lower energy requirements for regeneration. However, the low reaction rate of C02 with tertiary amines limits the use of MDEA solutions. Recent research (Chakravarty et al., 1985)indicates that the blends of primary and tertiary amines, mixtures of MEA and MDEA, can be used to enhance the absorption rate of C02 and the loading capacity, which brings about a considerable improvement in absorption and a great saving in energy requirement. The solubility data for COZ in the mixed amines at various temperatures and concentrations are required for acid gas purification process design. Many investigators determined the solubility of C02 in single amines and in mixed amines. Representation of experimental data with a thermodynamically rigorous model is also needed so that one can systematically correlate and predict the vaporliquid equilibrium (VLE) for these systems. For most of the single amine and mixed amines, due to the chemical reactions and neutral species that occur in the system, the modeling is difficult and complicated. The previous models investigated for single amine systems have been reviewed in detail by Austgen et al. (1989) and are not introduced here again. In this paper, we will focus on the modeling of mixed amine systems, for which some authors have made the exploration and obtained some progress. Chakravarty (1985)first extended the model of Deshmukh and Mather (1981) to systems including two amines (MDEA-MEA and MDEA-DEA). Activity coefficients were represented with Guggenheim's equation (Guggenheim, 1935). However, at that time there were no experimental acid gas solubility data reported in the literature for these mixed amine systems that could be t Permanent address: Department of Chemical Engineering, Tsinghua University, Beijing 100084,China.
used to validate the extension for the amine mixtures of interest. Austgen et al. (1991) used Chen's electrolyteNRTL equation (Chen and Evans, 1986) to correlate the experimental VLE data for MEA-MDEA-CO2-HzO and DEA-MDEA-CO2-HzO systems successfully. Recently, Li and Shen (1993) correlated their data for mixtures of MDEA and MEA using the Kent-Eisenberg model. Although they were able to reproduce their experimental results, the inconsistency of the data with the results of Jou et al. (1994a) makes the resulting correlation of questionable use for industrial calculations. In this paper, we use Pitzer's new excess Gibbs energy equation (Pitzer, 1991)to correlate the VLE data for MDEA-COZ-H~Oand MEA-C02-H20 systems and to predict those for the MDEA-MEA-COZ-H~O system.
Thermodynamic Framework Chemical Equilibria. In the aqueous phase for the MEA-MDEA-C02-H20 system, the following chemical equilibria are involved: C02 + 2H20
Ki
H 2 0 + MEAH'
Kz
+ MDEAH'
K3
H,O
RNHCOO- + H 2 0
K4
RNH,
+ HCO;
(la)
H30+ + MEA
(2a)
H30'
H30'
+ MDEA
(3a)
+ HCOL (RNH, = MEA) (4a)
The thermodynamic equilibrium constants used in this work are based on the mole fraction scale. The temperature dependence of the equilibrium constants is represented by the following function: 1nK = C, + C2/T+ C, In T + C,T
(1)
Coefficients Cl-Cd for reactions la-4a are taken from Austgen et al. and are summarized in Table 1. The Henry's constant has the unit of pascals. Its temperature dependence is expressed by the same functional form as shown in eq 1. Standard States. In this work, both water and alkanolamine are treated as solvents. The standard state associated with each solvent is the pure liquid at the system
0888-5885/94/2633-2006$04.50/00 1994 American Chemical Society
Ind. Eng. Chem. Res., Vol. 33, No. 8,1994 2007 Table 1. Temperature Dependence of the Equilibrium Constants for Reactions la-4a and Henry’s Constants for C02 In Hor In K = C ,
+ C 2 / T + C , In T + C,T
reaction compd c1 c2 c3 c4 temp range, OC -12092.1 -36.7816 0.0 0-225 coz 231.465 la MEA 2.1211 -8189.38 2a 0.0 -0.007484 +50 -9.4165 3a MDEA -4234.98 0.0 0.0 25-60 4a MEA 2.8898 -3635.09 0.0 0.0 25-120 -8477.7 11 -21.95743 0.005781 c02 170.7126 0-100 Henry’s const 0 Edwards et al. (1978). Bates and Pinching (1951). Schwabe et al. (1959). Austgen et al. (1989). e Chen et al. (1979).
temperature and pressure. The adopted standard state for ionic solutes is the ideal, infinitely dilute aqueous solution (infinitely dilute in solutes and alkanolamine) at the system temperature and pressure. The reference state chosen for molecular solute (C02) is also the ideal, infinitely dilute aqueous solution a t the system temperature and pressure. This leads to the following unsymmetric convention for normalization of activity coefficients: solvents:
fs+l,
ionic and neutral solutes:
asxB-l
fi*
- 1, xi
0, xszw = 0
gE=gS+$H
C
d e
Fa = 2xa/xI
(6)
Here the solvent-solvent interaction parameters Antnand Am! are defined as follows:
+ 3antnn Ann,= 2um, + 3a,, A,,,,, = 2a,,
The ion-solvent interaction parameter W,, below:
+
W,, = (2wDc 2w, - W,
where the subscript s refers to any nonaqueous solvent, i refers to ionic or neutral solutes, and w refers to water. Activity coefficients of all species are assumed to be independent of pressure. For simplicity, in this work, we neglect the molecular species of C02 and the ionic species of C032-in aqueous phase, because their concentrations are very low compared with the other species dissolved in mixed solvents. We also neglect the nonideality for gas phase, for our calculation is limited in the low-pressure range. Thermodynamic Expression. The original Pitzer (1973) general equation does not consider the solvent molecules in the system as interacting particles. Thus it is not suitable to the thermodynamic description of mixed solvent systems. In this paper, we use the new Pitzer equation (1991), in which all the species in a system are considered as interacting particles. In this model, the excess Gibbs energy is assumed to consist of short-range forces @) and long-range Debye-Huckel (gDH)terms:
source a b
(7) (8) is defined
+ 2unC+ 2una)/4
(9)
Here
Differentiation of eq 3 yields expressions for the shortrange force contribution to the activity coefficient product of salt MX and activity coefficient of neutral species N:
(2)
For the short-range term, only two- and three-suffix Margules expansions are used (Pitzer and Simonson, 1986): where the four terms containing W1,ijconvert the activity coefficient product from the pure fused salt to the infinite dilution reference state (subscript 1 denoting the solvent, water). The activity coefficient of a neutral species (N), on the same basis, is given by n
n‘
where subscripts c, a, and n and n’ represent cation, anion, and neutral species, respectively. The total mole fraction of ions (XI) is given by XI =
1-
cxn
(4)
n
The cation and anion fraction Fc and F a are defined for fully symmetrical systems by Fc = 2xc/xI
(5)
It should be pointed out that we only use a few parameters and delete many terms (e.g., quaternary terms) in the Clegg-Pitzer equations (see Clegg andPitzer, 1992). From eqs 12 and 13 it can be seen that only W,,, Antn,and A,. appear in the expressions. Therefore only these variables
2008 Ind. Eng. Chem. Res., Vol. 33, No. 8, 1994 Table 2. Temperature Dependence of the Density for Pure Solvents solvent MW expression H2O 18.02 d = 0.999382 + 0.00007208t - 7.28491 X 1 V t 2+ 2.65177 X lo-%* MEA 61.09 d 1.0325 - 0.0007979t MDEA 119.16 d = 1.0553 - 0.0007663t 0
temp range, "C 30-75 20-90 20-90
source a
b b
Littel et al. (1992). b Wang et al. (1992).
are treated as adjustable parameters and those in aij and U i j j are not. For long-range terms, the following expression, deduced from the original Pitzer equation (1973) with the transformation from a molality-based system to a mole fraction basis, is used (Clegg and Pitzer, 1992):
for mixed solvents is calculated as follows:
Here B, is similar to )0 : in the original Pitzer (1973) equation. The mole fraction ionic strength I, is defined as
Vn is the molar volume of the pure solvent n. The temperature dependences of the density and dielectric constant of the pure solvents are listed in Tables 2 and 3, respectively. The expressions of activity coefficient for ion M and solvent N are as follows:
I, = '/2xt3i i
where 4, is the volume fraction of solvent n:
The function of g ( x ) is expressed by
Here
I = '/2cz:ci
(18)
+ (1 - zM2/21,)exp(-d,'I2)
i
A, is the Debye-Huckel parameter on a mole fraction basis: A, = A,(cCn)'I2 n
Ci and C, are the molar concentrations of the ion i and solvent n, respectively; I is the ionic strength in molar concentration. A+is the original Debye-Hackel parameter (Pitzer, 1973), which is a function of temperature (Z?, density (d,) and dielectric constant (Dm) of the mixed solvents: 78.54 X 298.15 A, = 0.391( DmT
)
312
dm 'I2 (0.99702) (20)
where 0.391 is the A, value for pure water as solvent a t 25 "C. The density dm has units of g/mL. The parameter p is related to the hard-core collision diameter, or distance of closest approach, of ions in solution. For the NaC1H2O system over an extremely wide range of temperature and density, p was proposed (Pitzer and Li, 1983) as follows: p = 2150(d/DT)'I2
(21)
For mixed solvents, we use the following equation to calculate the p values: 1 8 . 0 2 d m ~ c ,112 p=2150(
1000DmT
)
(22)
According toRaatschenet al. (1987),the dielectric constant
The mixed solvent density data for the MEA-MDEA-
H20 ternary system are taken from the literature (Li and Shen, 1992). The mixed solvent density data for MJ3AH20 and MDEA-HzO binary systems are taken from the papers of Littel et al. (1992), Jou et al. (1986), and AlGhawas at al. (19891, respectively. Data Regression: Determining Interaction Parameters Adjustable interaction parameters are assumed to be temperature dependent and were fitted to the following simple function of temperature:
B (or W or A) = a + b/T
(27)
Binary Systems. For MEA-H20 and MDEA-H20 binary systems, we regressed the Margules interaction parameters (A12 and A21) by using the NRTL binary interaction parameters, which were given in Austgen's dissertation (Austgen, 1989). The interaction parameters for the MDEA-H20 system are assumed to be independent of temperature. The A12, A21, a, and b values thus obtained are listed in Table 4. Because there is no information for the MDEA-MEA binary system, we have to assume the interaction parameters between MDEA and MEA are equal to zero. Ternary Systems. For the MDEA-COpH20 system, there are only four main species (two neutral solvents,
Ind. Eng. Chem. Res., Vol. 33, No. 8,1994 2009 Table 3. Temperature Dependence of the Dielectric Constant for Pure Solvents
solvent
HzO MEA MDEA 4
expression 1 V ( t - 25) + 1.19 X 1 P ( t - 25)2- 2.8 X 10% 1 1 D = 35.76 + 14836(- - T 273.15)
D = 78.54[1- 4.579
- 25)3]
X
temp range, OC 0-100
source
-70 to 60
b
e)
D = 24.74 + 8989.3(~1 1
a
25-50
C
*
Maryott and Smith (1951). Ikada et al. (1968). Austgen et al. (1991).
Table 4. Fitted Values of Interaction Parameters between Solvents*
Am) = a + b/T Ad n-n' MEA-H20 H20-MEA MDEA-H20 H20-MDEA 4
A d
a
A31 A13
2.739 286 0.268 643 -0.169 097 -1.624 403
Azl A12
b, K -1318.671 -269.2729
25 O C -1.683 558 -0.634 503
40 OC -1.471 703 -0.591 242
80 OC -0.994 739 4 4 9 3 846
120 OC -0.6148 -0.4162
-0.169 097 -1.624 403
0 0
Subscripts 1 = H20, 2 = MDEA, 3 = MEA.
Table 5. Fitted Values of Interaction Parameters for MDEACOrH2O System.
B (or W) = a + b/T B (or W) a
BMX W1,MX w2,MX a
-232.080 3 -19.024 12 13.073 49
b, K 54909.07 5407.834 -4127.374
25 "C -47.914 35 4 8 8 6 156 -0.769 786
40 "C -56.735 97 -1.754 971 -0,106 687
80 OC -76.596 57 -3.710 984 1.386 183
120 "C -92.415 85 -5.268 979 2.575 278
Subscripts: 1 = H20,2 = MDEA, M = MDEAH+, X = HCOs-
MDEA and H20, and two ionic species, MDEAH+ and HC03-) in the equilibrated liquid phase. We approximately assume that all the dissolved C02 is converted into HC03- ions. The calculation of the concentration for each component is as follows: C~EAH =+ CHCO, = C'MDEA~
1000
-
I
I
0 Jou e t a1.(1994a)
100
a s v
10
0"
(28)
1
u
I
0
:
1
-a
0.1
2
From reactions l a and 3a, we get HK3xMfM*xXfX* pcOi = K,x,flx j 2
m 3
%P
0.01
GOO1
0 001
Here degree ( O ) represent the original concentration and a the COz loading in equilibrated liquid phase, expressed in mol of COz/mol of amine. In this work, we use our experimental VLE data (30 w t 7% MDEA at 25,40,80, and 120 "C, COPloading below 1.0 (Jou et al. 1994a)) to regress three interaction parameters ( W I , ~W2,MX, , and B m ) and their temperature coefficients (a and b). Here 1, 2, M, and X represent H20, MDEA, MDEAH+, and HC03-, respectively. The objective function 3 for regression is:
The parameters thus obtained are listed in Table 5, and the correlation results are shown in Figure 1. The average relative deviation of partial pressure of C02 for correlation is 7.32 5%. The large negative values of B m show that the MDEAH+ and HC03- ions are strongly associated into
I
I
1
0 01
01
1
CO, loading (mol CO,/mol MDEA)
Figure 1. Solubility of C02 in 30 w t % MDEA aqueous solutions at 25,40, 80, and 120 OC.
ion pairs in this solution. Accordingto the regular solution theory, if one neglects the three-particle interaction, the Margules parameter W i j or ajj is proportional to the difference of the attraction energy between the unlike molecules (i-j) and the like molecules (i-i, j-j). In our case, the more negative value is for W I , ~the , more hydration abilities of M+-H20 and X--H20 there are. The more negative values of W I , than ~ W2,m indicate the hydration abilities of ions by water are stronger than their solvation by amine. We use these regressed interaction parameters to predict the VLE data for 35 7% MDEA-COZ-H~Osystems at 40 "C (Austgen et al., 1991). The prediction results are shown in Figures 2 and 3. For the MEA-C02-H20 system, due to the formation of carbamate, there are five main species (two neutral
2010 Ind. Eng. Chem. Res., Vol. 33, No. 8, 1994 1000,
T a 100 Ir:
v
$
I
t
where 1, 3, N, X,and Y represent HzO, MEA, MEAH+, HCOs-, and RNHCOO-, respectively. Define
I
I
I
Jou e t a1.(1993) -prediction
x3xx K4flfY* (Kc)4= -= XlXY f3fx*
101
c
the [ value can be obtained by solving the quadratic equation:
0
z3
0.1
2a
t
i
'
0001
(39)
0 001
I
I
I
0 01
01
1
From equilibria la, 2a, and 4a, one gets
CO, loading (mol CO,/mol MDEA)
Figure 2. Solubility of COz in 35 w t % MDEA aqueous solutions at 40 and 100 "C. 1
"0
___________~_
r'
--
~
From equilibria l a and 2a, one gets
Austgen e t al (1991) -prediction
,id
10
-1
I
a
I
/
I I
'
@c0,)2 =
d'
HKfldN*XxfX* Klxlflx3f3
Actually, @col)1= (PCOJ2
(43)
The objective function (3)for regression is
J OB), and high MEA % in the amine mixture. These prediction curves are located in order (from low MEA % to high) and intersected at about a = 0.6. We also use our regressed parameters to predict the VLE data for the 2 M MDEA + 2M MEA aqueous system at 40 and 80 "C (from Austgen et al. 1991). The deviations are nearly in the same order as ours. See Figures 11and 12. We calculate the liquid-phase concentration of the CO2-loadedaqueous solutions with different MEA/MDEA ratios. The results are shown in Figures 13-16. It can be seen from these figures that the curve for the carbamate ion is sharper in high MEA mixture and seems flat in low MEA concentration. Because there is an appreciable amount of free C02 dissolved in the liquid phase above a
0 001 00
01
02
03
04
0 5
05
CO, loading (mol CO,/mol amine)
Figure 10. Solubility of COz in mixed MDEA and MEA aqueous solutions at 120 "C (MDEA + MEA = 30 wt %).
= 0.8, which we do not account for, our calculation is not valid in these regions.
Discussion There are several limitations which restrict the use of the model for mixed amine systems. The carbonate and free C02 species have not been considered; hence the model is expected to be unreliable a t very low and very high loadings. In our calculations, the MDEA-H20 parameters are temperature-independent. However, MDEA is known to have a high enthalpy of solution in water. This feature will invalidate any extrapolation of the model to very low loadings. At 120 "C, the partial pressures are in the range of lo00 kPa. Here gas-phase nonidealities may be significant and the use of fugacity coefficients calculated from an equation of state should be added. The equilibrium constants for MEA and MDEA are known only over a limited temperature range (see Table 1). Therefore the present calculations may be less accurate at higher temperatures where the equilibrium constants have been extrapolated. Nevertheless, from our computations, we have found the Clegg-Pitzer equations to be flexible and suitable for thermodynamic calculations of mixed aqueous amine
Ind. Eng. Chem. Res., Vol. 33, No. 8,1994 2013 Table 7. Average Deviation for Correlations and Prediction of VLE Data in Amine Systems and Sources of Experimental Data ___ ~~
6*#(%)
system MDEA-H20 MEA-HzO COz-MDEA-HzO
temp ("C)
amine concn
C02-MEA-H20
corr
0-5 M 0-5 M 30wt % 35wt % 2,4.28 M 30wt % 30wt % 2.5 M 0.5-3.4 M
25 45-160 25,40,80,120 40,100 40 25,40,80,120 0,60,100,150 40,W 40 25 40 80 120 40 80
pred
Figure
0.614 0.052 7.32
a
a
15.2 22.1 13.5 38.1 9.15 27.4 38.3 35.6 37.0 34.8 30.6 33.3
'
I
I
0 06
I
I
I
0 Austgen e t a1 (1991)
-prediction
\
_
d
92 06 06 CO, loading (mol CO,/mol
05
00
12
amine)
Figure 11. Solubility of C02 in 2 M MDEA + 2 M MEA aqueous solutions at 40 "C.
b
2 3 4 5
C
d e e
d
3ZMEA-27xMDkA a t 40°C
f
7 8 9 10
b b b b d d
11 12
et al. (1971). 8 6* =
HCO,
I /,
I
'=
0 00 ? _ .
1
6
MEAIMDEA = 1.5128.5 MEA/MDEA = 3/27 MEAIMDEA = 10120 MEAIMDEA = 20110 2 M MEA + 2 M MDEA 2 M MEA + 2 M MDEA 0 Austgen (1989). b Jou et al. (1994a). Jou et al. (1993). Austgen et al. (1991). e Jou et al. (1994b). f Murzin COTMDEA-MEA-H~O
source
02 0 4 c5 38 CO, loading (mol C02/mol amine)
Figure 13. Liquid-phase concentration of a COz-loaded 3 w t % MEA + 27 w t % MDEA aqueous solution at 40 OC.
10000
1 000
0 04
h
Q
a
6
100
N
0
E
u
5
0 003
10
Q)
5 I
aE
1
L
a
5
002
01
4 Li
6
a
00i
0 0 0 1 .I 00 CO,
I
1
I
I
I
0 2
0 4
06
08
10
loading (mol C02/mol amime)
\ 0 00 00
2
0 2
0 4
06
08
I I
10
CO, loading (mol C02/mol amine)
Figure 12. Solubility of COz in 2 M MDEA + 2 M MEA aqueous solutions at 80 OC.
Figure 14. Liquid-phase concentration of a CO2-loaded 10 w t % MEA + 20 w t % MDEA aqueous solution at 40 OC.
solutions with chemical equilibria. At first glance the above equations appear incompromisingly long and to contain very many parameters. However, they are linear equations in A, B, and W, which may be regressed without serious computationaleffort. To simplify the calculations,
many parameters are neglected. This model can still predict the vapor-liquid equilibria for mixed aqueous amine solutions using only the parameters obtained from single amine aqueous systems, without any additional adjustable parameters. Although the accuracy of the
2014
Ind. Eng. Chem. Res., Vol. 33, No. 8, 1994 L '9
I
I
20ZMEA- 10ZMDEA a t 40'C
OC
n ? "_ c 4 Ob 08 CO, loading (mol C02/mol amine)
10
Figure 15. Liquid-phase concentration of a Con-loaded 20 wt % MEA + 10 wt % MDEA aqueous solution at 40 OC.
'
30wtZMEA (40°C)
a, b = coefficients of eq 27 B = interaction parameter between ions C1, Cz, C3, C1= coefficients of eq 1 C = molar concentration D = dielectric constant d = density, g/mL F = ionic fraction 3 = objective function f = activity coefficient gE = excess Gibbs free energy H = Henry's constant, Pa I = ionic strength on molality basis I, = ionic strength on mole fraction basis K = thermodynamic chemical equilibrium constant,expressed in activity and in mole fraction K, = apparent chemical equilibrium constant, expressed in mole fraction concentration p = partial pressure, Pa or kPa as noted R = gas constant T = absolute temperature, K t = temperature, "C u = Margules interaction parameter V = molar volume W = interaction parameter between and among neutral and ionic species w = Margules interaction parameter x = liquid-phase mole fraction based on true species,molecular and ionic z = valence of an ion
Greek Letters a = Pitzer universal constant in eq 14 a = CO2 loading in liquid phase, mol of COz/mol of amine
6 = average relative deviation, 3'6
5 = molar concentration of carbamate ion p = C#J
hard-core collision diameter
= volume fraction
Superscripts O
OC
si
0 4
06
08
10
CO, loading (mol C02/mol MEA)
Figure 16. Liquid-phase concentration of a COz-loaded 30 wt % MEA aqueous solution at 40 OC.
prediction is still not high, it may be improved by adding , W 3 , ~parameters. y The mole the B w , WI,MY,W 2 , ~ yand fractions of MDEA and MEA in mixed aqueous amine solutions are relatively low compared with that of water. Neglecting the A23 and A32 parameters may not affect the accuracy of the prediction substantially. The regressed parameters can illustrate some physical properties (such as the association and solvation) of the ionic species in this complicated solution. Thus the Clegg-Pitzer model is likely to be of greatest use for aqueous solutions containing mixed organic solvents and inorganic (or organic) salts. Acknowledgment This work was supported financially by the Natural Sciences and Engineering Research Council of Canada under Strategic Grant 100777. Nomenclature A = interaction parameter between and among neutral molecules A4 = Debye-Hiickel parameter for osmotic coefficient A, = Debye-Htickel parameter on mole fraction basis a = Margules interaction parameter
= initial concentration
* = unsymmetric convention
DH = long-range term S = short-range term Subscripts
a, X, Y = anion c, M, N = cation I = ionic species i, j = any species m = mixed solvent system N, n, n', s = neutral solvent species w = water 1 = HzO 2 = MDEA 3 = MEA M = MDEAH+ N = MEAH+ X = HC03Y = RNHCOOcal = calculated value corr = correlation exp = experimental value pred = prediction Literature Cited Al-Ghawas, H. A.; Hagewiesche, D. P.; Ruiz-Ibanez, G.;Sandall, 0. C. Physicochemical Properties Important for Carbon Dioxide Absorptionin Aqueous Methyldiethanolamine.J. Chem.Eng.Data 1989,34,385.
Austgen, D. M. ModelingPhaeeEquilibriainAcid Gas-AlkanolamineWater Systems.Ph.D. Dissertation, The Universityof Texas, 1989.
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Received for review April 20, 1994 Accepted May 10,1994 *
* Abstractpublished in Advance ACSAbstracts, June 15,1994.
