Use of Flow Calorimetry for Determining Enthalpies of Absorption and

Experimental data served to obtain the integral enthalpies of absorption and for indirect determination of solubility limits. .... Dissolution Mechani...
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Ind. Eng. Chem. Res. 1998, 37, 4136-4141

Use of Flow Calorimetry for Determining Enthalpies of Absorption and the Solubility of CO2 in Aqueous Monoethanolamine Solutions C. Mathonat,†,‡ V. Majer,*,† A. E. Mather,*,§ and J.-P. E. Grolier† Laboratoire de Thermodynamique et Ge´ nie Chimique, Universite´ Blaise Pascal/CNRS, 63177 Aubie` re, France, and Department of Chemical & Materials Engineering, University of Alberta, Edmonton, Alberta T6G 2G6, Canada

A flow mixing unit adapted to a Setaram C-80 calorimeter was used for measuring enthalpies of absorption of carbon dioxide in a 30 wt % aqueous solution of monoethanolamine (MEA) at three temperatures (313.15, 353.15, and 393.15 K) and three pressures (2.0, 5.0, and 10.0 MPa). Determinations were performed both in the region where the gas is fully absorbed in the solvent and also in the region of concentrations above the saturation. Experimental data served to obtain the integral enthalpies of absorption and for indirect determination of solubility limits. Where comparison was possible, the presented results derived from calorimetric determinations were in reasonable agreement with those obtained from phase equilibria measurements. Introduction Gaseous effluents encountered in different industrial processes contain considerable amounts of carbon dioxide, which is of major concern considering its negative impact on the environment (carbon dioxide is mainly involved in the greenhouse effect, which refers to a net warming of the earth’s atmosphere). Absorption (physical and/or chemical) in aqueous solutions of polar organic substances is frequently used for removal of CO2 and other acid gases (such as H2S) from gaseous streams and their subsequent recovery by desorption. In particular, aqueous alkanolamines are useful for this purpose and different studies have been undertaken to obtain thermodynamic and kinetic data describing this process and have been reviewed by Astarita et al. (1983) and Kohl and Riesenfeld (1985). Monoethanolamine (MEA) and diethanolamine (DEA) are among the most widely used alkanolamines in industrial acid gas removal of CO2 or H2S from gas streams. Chemical absorption of CO2 in aqueous solutions of primary and secondary alkanolamines is controlled by the chemical reactions that take place in the solution (Caplow, 1968; Danckwerts, 1979; Blauwhoff et al., 1984; Glasscock et al., 1991). The primary or secondary alkanolamine RR′NH (where R′ is a hydrogen atom or an alkyl group) reacts directly with CO2 to form an intermediate zwitterion:

CO2 + RR′NH a RR′NH+CO-

(1)

The zwitterion reacts subsequently with a molecule of amine to form a carbamate by the intermediate of hydrogen attached to the nitrogen of the amine: * To whom correspondence should be addressed. V.M.: Fax (33) 473 40 71 85; E-mail [email protected]. A.E.M.: Fax+1(403)492-2881;[email protected]. † Universite ´ Blaise Pascal/CNRS. ‡ Present address: Setaram, 7 rue de l’Oratoire, 69300 Caluire, France. § University of Alberta.

RR′NH+COO- + RR′NH a RR′NH2+ + RR′NCOO(2) It is apparent from the previous two equations that reactions between CO2 and primary or secondary alkanolamines are stoichiometrically limited to 0.5 mol of CO2 per mol of amine. However, absorption of CO2 exceeding this limit is possible considering physical absorption and hydrolysis of the zwitterion:

RR′NH+COO- + H2O a H3O+ + RR′NCOO-

(3)

Information on solubilities and heats of dissolution (reaction) of acid gases (mainly CO2 and H2S) in aqueous solutions of alkanolamines is of primary importance for designing unit operations of acid gas removal. It is also necessary for a complete thermodynamic description of these systems allowing the prediction of phase equilibria and heat effects outside conditions where experimental data are available. The most systematic studies of phase equilibria between CO2 and different aqueous solutions of alkanolamines have been undertaken at the University of Alberta. Equipment and procedures used in the experiments have been described in detail by Jou et al. (1994). Solubility data were obtained by using a closed system where the gas was circulated and bubbled through an amine solution contained in a windowed equilibrium cell, which was a Jerguson gauge. Gas was added in amounts monitored by the pressure in the cell, and if necessary, nitrogen was added to maintain the total pressure above atmospheric pressure. A magnetic pump was used to circulate the vapor phase in the solvent for at least 8 h to ensure that equilibrium was reached. Sampling of both phases were performed; samples of vapor were analyzed by gas chromatography, and liquid samples were first injected into NaOH solution and then analyzed by titration. By using the direct measurements of solubility, an attempt was made to determine quantities related to the energy evolved during the absorption and estimates of the differential enthalpy of gas in the solution were obtained. The calculation procedure has been described by Lee et al. (1973). An apparatus similar to that described above

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Ind. Eng. Chem. Res., Vol. 37, No. 10, 1998 4137

was used recently by Jou et al. (1994) for measurements of absorption of CO2 in aqueous solutions of monoethanolamine, methyldiethanolamine, and their mixtures of different concentrations. Results were obtained at temperatures of 298.15, 313.15, 353.15, and 393.15 K over a range of pressures from 100 kPa to 20 MPa at partial pressures of CO2 ranging from 0.001 to 19 930 kPa. Flow calorimetry offers an inverse approach to investigating absorption of gases in solutions compared to phase equilibria measurements. It makes it possible to measure directly the enthalpy of absorption of a gas in an aqueous solution of an alkanolamine and to derive from the shape of the calorimetric curve the solubility of a gas in a solution. Such direct calorimetric measurements have been undertaken at Brigham Young University. The results have been obtained concerning the absorption of CO2 in diglycolamine (DGA) (Christensen et al., 1986), methyldiethanolamine (MDEA) (Merkley et al., 1987), and diethanolamine (DEA) (Oscarson et al., 1989) solutions. All the data were obtained on instruments described by Christensen et al. (1976, 1981), which were essentially isothermal heat leak flow calorimeters working in the temperature ranges between 273 and 343 K and 273 and 423 K, respectively, and at pressures from ambient to 41 MPa. The measurements reported for absorption of CO2 in solutions of alkanolamines were performed in the range of temperatures between 288.7 and 422.0 K and pressures of 0.087-1.466 MPa with an estimated accuracy of (5% and a precision of (4%. Recently, we have reported new results obtained for CO2 absorption in the MDEA solutions (Mathonat et al., 1997) over the temperature range from 313.15 to 393.15 K at pressures 2.0, 5.0, and 10.0 MPa which is higher than those used in similar measurements at the Brigham Young University (see above). The calorimetric measurements were used to determine two important thermodynamic characteristics, which are the limit of solubility of the gas in the solvent and the enthalpy of absorption reflecting both the heat effect due to the physical dissolution of gas in the solvent and to the chemical reaction between CO2 and MDEA. The direct calorimetric results were converted into differential enthalpies and compared with those obtained by calculation from the solubility measurements made at the University of Alberta (Jou et al., 1982). This paper is analogous to the previous one and we present here results obtained in a study of CO2 absorption in a 30 wt % MEA solution. No calorimetric measurements have been undertaken before now on absorption in the aqueous solution of this amine over the range of temperatures and pressures studied. Experimental Section The equipment and experimental procedures used in this study are similar to those used in our previous investigations (Mathonat et al., 1994, 1997); therefore only the most salient information will be given here. The measurements were carried out in a differential Setaram C80 heat conduction calorimeter using a flow mixing unit of our own construction. The differential thermal signal from the thermopiles surrounding the cells is proportional to the heat effect measured and can be related to the amount of energy after electrical calibration. The mixing cell is made of a stainless steel tube coiled in about 45 loops (2.4 m length) inside a

Figure 1. Measured enthalpies ∆Habs (kilojoules per mole of MEA) versus CO2 loading R (moles of CO2 per mole of MEA) for absorption of CO2 in an aqueous solution 30 wt % MEA at 313.15 K and 2.0 MPa. Table 1. Limit of Solubility of CO2 in Aqueous Solution of MEA (30 wt %) Obtained from Calorimetric Measurements Compared with the Values from the Literature R (mol of CO2/mol of MEA) T (K)

p (MPa)

this work

Jou et al. (1994)a

313.15 313.15 313.15 353.15 353.15 353.15 393.15 393.15 393.15

2.0 5.0 10.0 2.0 5.0 10.0 2.0 5.0 10.0

0.90 1.07 0.95 0.79 0.89 0.99 0.55 0.72 0.88

0.92 1.03 1.09 0.73 0.86 0.95 0.61 0.70 0.78

a Literature values obtained by interpolation of the experimental data.

metallic confinement cylinder that fits into the well surrounded by the thermopile detector in the calorimetric block. The reactants were supplied to the system from two metering pumps. A Gilson Master-305 pump operated in the range between 100 and 500 µL‚min-1 and was used for the aqueous solution of alkanolamine. An Isco-314 pump supplied the gas at flow rates from 10 to 500 µL‚min-1. In the last period of our work we have used also a new Isco 260 pump, which allowed us to work without calibration and with a better control of the flow rate, especially at low values. During measurements with low concentrations of carbon dioxide at the highest experimental pressure (10.0 MPa), the flow rates of CO2 were very low (down to 10 µL‚min-1); therefore, we restarted the measurements specifically for this pressure using this new pump to check the results obtained previously. All the flow circuit serving to introduce the fluids from pumps to the mixing cell through a preheating system and bringing the mixture outside the calorimeter cell across a back pressure controller is made of stainless steel capillary tubing (o.d. 1.6 mm, wall thickness 0.3 mm). Before entering the cell, fluids passed first through a heated countercurrent heat exchanger made of the two inlet tubes and the outlet tube coiled tightly

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Figure 2. (a) Measured enthalpies ∆Habs (kilojoules per mole of CO2) versus number of mole of solution (water + MEA) per mole of CO2 for absorption of CO2 in an aqueous solution of 30 wt % MEA at 313.15 K. (0) 2.0 MPa, (2) 5.0 MPa, (]) 10.0 MPa. (b) Measured enthalpies ∆Habs (kilojoules per mole of CO2) versus number of moles of solution (water + MEA) per mole of CO2 for absorption of CO2 in an aqueous solution of 30 wt % MEA solution at 353.15 K. (0) 2.0 MPa, (2) 5.0 MPa, (]) 10.0 MPa. (c) Measured enthalpies ∆Habs (kilojoules per mole of CO2) versus number of moles of solution (water + MEA) per mole of CO2 for absorption of CO2 in an aqueous solution of 30 wt % MEA solution at 393.15 K. (0) 2.0 MPa, (2) 5.0 MPa, (]) 10.0 MPa.

together and covered with a heating tape. The temperature of the fluids was further thermoregulated during their passage through a temperature-controlled head located just above the calorimeter block. The temperature of the head is controlled within (0.01 °C and adjusted such that the circulation of a fluid through the mixing cell does not induce any heating effect. The adjustment of the required pressure and maintenance of its constancy to better than 0.02 MPa is assured by means of a back pressure regulator. A Grove-Mity-Mite Model 91 pressure controller with an operating range up to 20 MPa was connected at the end of the flow system. The controller has a dome filled with nitrogen determining the experimental pressure, which is separated from the flow circuit by a membrane commanding a seat opening toward the ambient pressure side.

During an experiment, the aqueous solution was injected in the system usually first at a flow rate corresponding to a specified concentration. When a baseline is obtained, the calorimetric signal is recorded over 15-30 min. Then the second pump with CO2 is started and the thermal mixing effect is observed. A steady state is attained after about 20-30 min from the beginning of mixing and the signal is recorded for about 1 h. Once the plateau characterizing mixing was well established, the injection of amine was stopped in order to record the baseline for CO2. The signal usually did not differ significantly from that obtained with the amine. Before mixing was started, the pressure of CO2 in the Isco pump had been adjusted at a value slightly below the system pressure in order to avoid perturbation of the calorimeter signal when CO2 is injected. The use

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Figure 3. Measured enthalpies ∆Habs (kilojoules per mole of CO2) when R tends to zero versus temperature. (b) This work for 30 wt % MEA. Table 2. Enthalpies of Solution of CO2 in Aqueous Solution of MEA below the Saturation Loading Point When r Tends to Zero Compared with Enthalpies of Absorption of CO2 in Aqueous Solution of MDEA in the Same Conditionsa this work wt % 30

a

T (K)

p (MPa)

313.15 2.0, 5.0, 10.0 353.15 2.0, 5.0, 10.0 393.15 2.0, 5.0, 10.0

Mathonat et al. (1997)

-∆Habs σ -∆Habs σ (kJ/mol (kJ/mol (kJ/mol (kJ/mol of CO2) of CO2) of CO2) of CO2) 81 90 102

3 5 3

49 55 58

4 3 5

Mathonat et al. (1997).

of two manometers allowed continuous comparison of the system and pump pressures. The first manometer was a Druck DPI 600 model with an uncertainty of 0.1% of the full scale (0-25 MPa), which was placed near the exit of the Isco pump in the CO2 line. It served for a precise pressure indication during an experiment and to adjust the pressure in the pump filled with CO2 before injection. The second manometer was a 0-10 MPa gauge, Besanc¸ on Instruments, precision 0.5% of full scale (0-10 MPa), located between the calorimeter and the back pressure regulator and serving for approximate pressure indication outside the periods when mixing is performed. Mixing of CO2 with aqueous solutions of alkanolamine implied several problems. The main difficulty occurred at the highest pressure where the flow rates of CO2 are small (typically between 10 and 20 µL‚min-1) and are very sensitive to possible fluctuations in the laboratory temperature. This is mainly due to the expansion or contraction of the fluid in the pump, which can lead to important changes in the flow rate when the temperature change is fast. The effect of temperature fluctuations can be documented with an example of measurements with CO2 in a pump containing 60 cm3 of CO2 supplied to the calorimeter at a flow rate of 10 µL‚min-1. A temperature increase of 1 K in 1 h will be accompanied with a 50% and 100% change in flow rate at a pressure of 2 and 10 MPa, respectively. This change

will be increasing (decreasing) proportionally with the increase (decrease) of the volume of CO2 contained in the pump. It is therefore advantageous to work with a pump only partially filled and cover the pump cylinder containing CO2 with a thermal insulation attenuating temperature fluctuations in the laboratory. A much less, yet not still negligible, effect is the instability in room temperature on the density of CO2 used in the conversion of the volumetric flow rate of gas to the number of moles used in the calculation of enthalpies of absorption per mole of CO2. The change of 1 K in the room temperature will cause an error of 0.5% and 1% in ∆Habs at a pressure of 2 and 10 MPa, respectively. In addition, fluctuations of room temperature can affect the pressure of nitrogen in the dome of the regulator, inducing pressure changes in the system that will affect ∆Habs when appropriate correction is not taken into consideration. An error in pressure of 0.1 MPa will be reflected by a change in heat of absorption related to 1 mole of CO2 of about 6% and 0.2% at a pressure of 2 and 10 MPa, respectively. The three sources of errors mentioned above are random in character and change between individual experiments; they are reflected in reproducibility of measurements. Measurements were restarted at different periods of time and the reproducibility of the data was about (3%. In the present study, the data representing certain CO2 loadings were taken at different flow rates (typically 100 and 200 mg‚min-1 for MEA increasing correspondingly the CO2 flow rates to keep the loading constant) to see if the results for ∆Habs can be affected when flow rate increases. We have observed no flow rate dependence in the range of our experimental flow rates. Thus we have assumed that equilibrium is reached under the various flow rate conditions. As mentioned above, measurements at 10.0 MPa were performed first with the Isco-314 pump and then with the new Isco-260 pump; we have observed no difference between the two sets of results. After analyzing possible sources of systematic and random errors and after taking into account fluctuations of room temperature in the laboratory, we have come to a conclusion that expected uncertainty of our calorimetric measurements is (4% at 2.0 and 5.0 MPa and (7% at 10.0 MPa. The chemicals employed, carbon dioxide (Alphagaz, N48, 99.998 mol % pure) and MEA (Fluka, g99% pure), were used without any further purification. Water was deionized and degassed before use. Results and Discussion One way to present the results of the measured enthalpy ∆Habs for CO2 in aqueous MEA-water solutions is shown in Figure 1 where ∆Habs is expressed as enthalpy of absorption per mole of MEA plotted against loading R (moles of CO2 per mole of MEA). This representation is mainly used to determine the limit of solubility of the gas in the solvent. ∆Habs increases with the CO2 loading first as a straight line that corresponds to the region where CO2 is completely absorbed in the aqueous solution of monoethanolamine; then ∆Habs grows slower to become quite constant after the saturation point (loading) was achieved and no additional CO2 is absorbed in the aqueous solution of alkanolamine. The limit between ranges where CO2 is completely absorbed and where the aqueous solution of alkanolamine is saturated is marked by a saturation loading point that

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Figure 4. Dependence of ∆Hdiff on CO2 loading. (s) Equation 5.

is dependent on temperature and amine solution composition. The saturation point can be determined easily from ∆Habs (kilojoules per mole of MEA) versus R curves as an intersection between a polynomial function describing data just before the saturation point (identified by a dashed line) and the straight line after this point. More than four parameters for the polynomial function are usually necessary for good description of the dashed curve. The saturation loading points determined by means of calorimetric measurements are reported in Table 1 and are compared with values of solubility limit of CO2 obtained by direct measurements in aqueous solutions of MEA (Jou et al., 1994). Comparison between the two kinds of measurements give a reasonable agreement (within 7%). A second way to present our results is to plot ∆Habs expressed per mole of CO2 versus the number of moles of solvent n1 (i.e., number of moles of water, nw + number of moles of alkanolamine, na) per mole of CO2, n2. This representation is useful to determine the ∆Habs value per mole of CO2 from the plateau of enthalpy of absorption obtained at elevated values of n1/n2, which corresponds to the region where CO2 is completely absorbed and where in the limit the value of R tends to zero. The results are presented in Figure 2 for each temperature and various pressures. In a first region that corresponds to the saturation of the alkanolamine solution (CO2 is not completely absorbed), ∆Habs can be described by a straight increasing line that passes through the origin. Finally the saturation loading point is attained, and in a second main region where CO2 is completely absorbed (solution is not saturated) the ∆Habs data do not change significantly with n1/n2. The values are, however, relatively scattered in this region, and at a given temperature ∆Habs seem to be independent of pressure. Therefore an average value was calculated, taking into account all data points without respect to pressure. The enthalpy of absorption per mole of CO2 can be also determined from the first representation (see Figure 1) as a slope of the linear ascending relationship of ∆Habs per mole of MEA versus loading R. According to our experience this method is, however, less convenient and less precise as the slope must be determined by a least-squares linear regression forcing the line

describing the experimental data to pass through the origin. However, due to errors in measurement the experimental points do not behave quite that way and results of fitting differ depending on the number of experimental data points considered in a correlation. The number of data points included in the fitting is always somewhat arbitrary as it is not usually clear at which value of R the shape of the curve is changing due to the proximity of the saturation point. For that reason we have preferred to associate ∆Habs per mole of CO2 with the average value of a plateau in Figure 2. In Figure 3 are plotted the average values of the enthalpy corresponding to the complete absorption region (unsaturated solution, R in the limit going to zero) as a function of temperature. It is apparent that absolute value of ∆Habs increases approximately linearly with temperature; the error bars correspond to the standard deviation calculated for each averaged value (data at all pressures treated together). The results are reported in Table 2 and are compared with those obtained for MDEA solutions in the same ranges of temperatures and pressures. It is apparent that ∆Habs values are more exothermic in the case of MEA than for MDEA. Reactions between a primary or secondary amine such as MEA leading to formation of a carbamate can be assumed as much more exothermic than is the case of reactions with a tertiary amine where carbamate is not formed. No calorimetric measurements of the enthalpy of solution of CO2 in MEA solutions have been reported. All values have been obtained by application of the Gibbs-Helmholtz equation

[

]

∂ ln fCO2 ∂(1/T)

x

)

H h 2 - H2• ∆Hdiff ) R R

(4)

to the experimental solubility data. Usually the approximation that the fugacity of CO2 is equal to its partial pressure was made. The values obtained from the solubility measurements are reliable to about (20% (Lee et al., 1974). This type of calculation does not allow any temperature dependence to be seen, as the procedure involves the differentiation of the partial pressure versus the reciprocal temperature over a narrow range of temperature. However, the calculated values indicate that the enthalpy is constant as R tends to zero. Murzin and Leites (1971) reported values of ∆Hdiff obtained from their own solubility data (Murzin et al., 1969), together with those of Mason and Dodge (1936) and Lyudkovskaya and Leibush (1949). Other values of ∆Hdiff were reported by Lee et al. (1974) and by Jou et al. (1994). The data are plotted in Figure 4 and the agreement is within (20%. The accuracy of ∆Hdiff obtained by that procedure is inherently inaccurate because of the process of differentiation applied to the experimental solubility data. Pearce (1978) reported a value of 83.9 kJ/mol of CO2 without explanation of how this number was obtained. The ∆Hdiff values from the three sources were fitted by a modified Gompertz equation [Davis (1943)]:

-∆Hdiff ) 83 - bcd

-R

(5)

where b ) 61.117, c ) 2.569 × 10-6, and d ) 134.78. The standard error of fit is (4.1 kJ/mol. To obtain ∆Hint for comparison with the present work, the equation was integrated with respect to R:

Ind. Eng. Chem. Res., Vol. 37, No. 10, 1998 4141

Literature Cited

Figure 5. Comparison of ∆Hint calculated from ∆Hdiff with experimental measurements at 353.15 K and 5.0 MPa. (s) Equation 7.

∫0R∆Hdiff dR

-∆Hint ) -

) 83R -

b Ei[dR ln c] b Ei[ln c] + ln d ln d

(6) (7)

The equation is plotted in Figure 5 for comparison with the experimental values obtained at 353.15 K and 5 MPa. In light of the inaccuracies of the ∆Hdiff, the agreement with the present work is very good. Conclusions Enthalpies of absorption of CO2 in an aqueous solution of MEA (30 wt %) have been measured in a range of temperatures and pressures that correspond to strong exothermic effects. Calorimetric measurements present advantages: they give with a reasonable accuracy both the solubility limit of the gas in the solvent (indirect determination) and the integral enthalpy of absorption of the gas completely dissolved in solvent (direct determination). The solubility of CO2 in aqueous solution of MEA decreases with increasing temperature and increases with increasing pressure. Analysis of the curves representing enthalpy of absorption as a function of concentration shows that in the region where the gas is completely absorbed in the aqueous solution of alkanolamine, ∆Habs increases approximately linearly with temperature between -81 and -102 kJ/mol of CO2 and is independent of the pressure within the experimental errors. Integral enthalpies of absorption obtained by integration of differential enthalpies of absorption are in agreement within the error bounds of both types of data.

Astarita, G.; Savage, D. W.; Bisio, A. Gas Treating with Chemical Solvents; Wiley: New York, 1983. Blauwhoff, P. M. M.; Versteeg, G. F.; Van Swaaij, W. P. M. A Study on the Reaction between CO2 and Alkanolamines in Aqueous Solutions. Chem. Eng. Sci. 1984, 39, 207. Caplow, M. Kinetics of Carbamate Formation and Breakdown. J. Am. Chem. Soc. 1968, 90, 6795. Christensen, J. J.; Hansen, L. D.; Eatough, D. J.; Izatt, R. M.; Hart, R. M. Isothermal High-Pressure Flow Calorimeter. Rev. Sci. Instr. 1976, 47, 730. Christensen, J. J.; Hansen, L. D.; Izatt, R. M.; Eatough, D. J.; Hart, R. M. Isothermal, Isobaric, Elevated Temperature, High-Pressure Flow Calorimeter. Rev. Sci. Instrum. 1981, 52, 1226. Christensen, S. P.; Christensen, J. J.; Izatt, R. M. Enthalpies of Solution of Carbon Dioxide in Aqueous Diglycolamine Solutions. Thermochim. Acta 1986, 106, 241. Danckwerts, P. V. The Reaction of CO2 with Ethanolamines. Chem. Eng. Sci. 1979, 34, 443. Davis, D. S. Empirical Equations and Nomography; McGrawHill: New York, 1943; p 57. Glasscock, D. A.; Critchfield, J. E.; Rochelle, G. T. CO2 Absorption/ Desorption in Mixtures of Methyldiethanolamine with Monoethanolamine or Diethanolamine. Chem. Eng. Sci. 1991, 46, 2829. Jou, F.-Y.; Mather, A. E.; Otto, F. D. Solubility of H2S and CO2 in Aqueous Methyldiethanolamine Solutions. Ind. Eng. Chem. Process Des. Dev. 1982, 21, 539. 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. Kohl, A. L.; Riesenfeld, F. C. Gas Purification, 4th ed.; Gulf Publishing Company: Houston, TX, 1985. Lee, J. I.; Otto, F. D.; Mather, A. E. Design Data for Diethanolamine Acid Gas Treating Systems. Gas Process./Canada 1973, 65 (4), 26. Lee, J. I.; Otto, F. D.; Mather, A. E. The Solubility of H2S and CO2 in Aqueous Monoethanolamine Solutions. Can. J. Chem. Eng. 1974, 52, 803. Lyudkovskaya, M. A.; Leibush, A. G. Solubility of Carbon Dioxide in Solutions of Ethanolamines under Pressure. Zh. Priklad. Khim. 1949, 22, 558. Mason, J. W.; Dodge, B. F. Equilibrium Absorption of Carbon Dioxide by Solutions of the Ethanolamines. Trans. AIChE 1936, 32, 27. Mathonat, C.; Hynek, V.; Majer, V.; Grolier, J.-P. E. Measurements of Excess Enthalpies at High Temperature and Pressure using a New Type of Mixing Unit. J. Sol. Chem. 1994, 23, 1161. Mathonat, C.; Majer, V.; Mather, A. E.; Grolier, J.-P. E. Enthalpies of Absorption and Solubility of CO2 in Aqueous Solutions of Methyldiethanolamine. Fluid Phase Equil. 1997, 140, 171. Merkley, K. E.; Christensen, J. J.; Izatt, R. M. Enthalpies of Absorption of Carbon Dioxide in Aqueous Methyldiethanolamine Solutions. Thermochim. Acta 1987, 121, 437. Murzin, V. I.; Leites, I. L.; Tyurina, L. S.; Perstina, Z. I. The Partial Pressure of CO2 above Slightly Saturated Water Solutions of Monoethanolamine. Khim. Prom. (Moscow) 1969, 785. Murzin, V. I.; Leites, I. L. Partial Pressure of Carbon Dioxide over Its Dilute Solutions in Aqueous Aminoethanol. Zh. Fiz. Khim. 1971, 45, 417. Oscarson, J. L.; Van Dam, R. H.; Christensen, J. J.; Izatt, R. M. Enthalpies of Absorption of Carbon Dioxide in Aqueous Diethanolamine Solutions. Thermochim. Acta 1989, 146, 107. Pearce, R. L. Hydrogen Sulfide Removal with Methyl Diethanolamine. Proc. Gas Process. Assoc. 1978, 57, 139.

Received for review November 3, 1997 Revised manuscript received June 2, 1998 Accepted June 15, 1998 IE9707679