Article pubs.acs.org/IECR
Assessing Anhydrous Tertiary Alkanolamines for High-Pressure Gas Purifications Paul M. Mathias,*,† Louis V. Jasperson,‡ David VonNiederhausern,‡ Mark D. Bearden,§ Phillip K. Koech,§ Charles J. Freeman,§ and David J. Heldebrant§ †
Fluor Corporation, 3 Polaris Way, Aliso Viejo, California 92698, United States Wiltec Research Company, 488 South 500 West, Provo, Utah 84601, United States § Pacific Northwest National Laboratory, Richland, Washington 99354, United States ‡
ABSTRACT: Anhydrous tertiary alkanolamines chemically react with CO2 and H2S, with greater selectivity for the latter. This is in direct contrast to aqueous amine-based solvent systems, which exhibit higher selectivity for CO2 over H2S. Anhydrous tertiary alkanolamines exhibit pressure-induced chemical fixation of CO2 to form zwitterionic ammonium alkylcarbonate ionic liquids, while the same tertiary alkanolamines react with H2S at atmospheric pressures to form hydrosulfide ionic liquids. This difference in capture pressure implies that certain anhydrous alkanolamines could be chemically selective for H2S over CO2. We present here the first published vapor−liquid−liquid equilibrium (VLLE) data of anhydrous ethyldiethanolamine (EDEA) with CH4, C3H8, H2S, and CO2 at 10−50 °C measured by the TPx and TPxy methods. The data are modeled in Aspen Plus using an NRTL-with-solvation model. Data trends and the underlying phenomena are discussed for each gas. We also present process simulations that compare anhydrous EDEA’s performance for CO2 and H2S high-pressure separations to other solvents such as Fluor Solvent (propylene carbonate), Selexol, and aqueous methyldiethanolamine (MDEA) for a representative gas-purification absorber. This work indicates that a niche for anhydrous EDEA in high-pressure gas purifications may be its stronger absorption for CO2 and H2S (relative to physical solvents) and its selectivity for H2S over CO2 (relative to chemical solvents).
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INTRODUCTION Chemical solvents such as alkanolamines are well-suited for gastreating applications, such as natural gas purification and CO2 capture, because they provide high selectivity for acid-gas components (e.g., CO2 and H2S) over nonpolar species.1,2 More than 80 years ago Bottoms3 first recognized amines as chemical solvents, and also invented the absorption-stripping process used today. Amine scrubbing is widely expected to be the dominant technology for the low-pressure CO2 capture from power plants, particularly for coal-fired plants.4 Amine scrubbing is also widely used for higher-pressure gas purifications associated with gas processing.1,2 However, chemical solvents have the disadvantage of relatively high enthalpies of solution, and in cases where the acid-gas impurities are at a significant composition of the total gas stream, the energy penalty of regenerating the solvent by application of heat may be out of proportion to the value of the treated gas. This has provided an opportunity for physical solvents, in which nonreactive organic agents are used. Here, solvent regeneration can often be accomplished by a simple reduction in pressure, without the need for heat to increase the temperature. Examples of physical solvents for gas-purification processes are Rectisol (methanol), Fluor Solvent (propylene carbonate), and Selexol (mixture of propylene glycol dimethyl ethers).1 In this work, we present a preliminary evaluation of the niche for anhydrous tertiary alkanolamines such as EDEA (ethyldiethanolamine) for high-pressure gas purifications involving H2S and CO2. These applications prefer high absorption capacity for both H2S and CO2, selectivity for H2S over CO2, and low regeneration costs. Chemical solvents usually have a thermodynamic selectivity for CO2 over H2S, and one option to mitigate © 2013 American Chemical Society
this undesired selectivity is to reduce the contact time since the kinetics of H2S absorption is much faster than that of CO2.5 Chemical solvents are typically aqueous solutions, and there is an energy penalty associated with the vaporization of water. Physical solvents usually have stronger absorption for H2S over CO2, but the absorption strength is low relative to chemical solvents. Researchers have adopted mixed-solvent processes that combine the bulk removal capability of the physical solvents with the strength of chemical solvents to achieve low residual acid-gas specifications, but these hybrid solvents tend to compromise the H2S selectivity of the physical solvents.1 Solvents like EDEA may have a special niche because they are water free, and have chemical absorption strength, excellent bulk capacity, and H2S selectivity. Anhydrous tertiary alkanolamines have been shown to chemically react with CO2 and H2S, but with selectivity for the latter. Anhydrous tertiary alkanolamines exhibit pressureinduced chemical fixation of CO2 under elevated pressure to form zwitterionic ammonium alkylcarbonate ionic liquids (Figure 1). The same tertiary alkanolamines react with H2S at atmospheric pressures to form hydrosulfide ionic liquids (Figure 1).6,7 This difference in absorption characteristics indicates that anhydrous alkanolamines can be chemically selective for H2S over CO2. Tertiary alkanolamines were chosen for this study because they are the only alkanolamines that do not directly react with Received: Revised: Accepted: Published: 17562
July 3, 2013 November 11, 2013 November 14, 2013 November 14, 2013 dx.doi.org/10.1021/ie4020974 | Ind. Eng. Chem. Res. 2013, 52, 17562−17572
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Figure 1. Anhydrous ethyldiethanolamine reacting with either H2S or CO2.
CO2 through the carbamate mechanism, thus they operate (at least with respect to CO2) closer to a pressure swing; as explained by Kohl and Nielsen,1 the regeneration of the rich solvent may be accomplished by pressure reduction (pressure swing) or temperature increase (temperature swing). Of the anhydrous tertiary alkanolamines, dialkanolamines were chosen due to their low vapor pressures and high polarities. A low vapor pressure is clearly a requirement for gas-separation solvents to reduce solvent slip. Less obvious is the need for a polar sorbent. Anhydrous chemically selective solvent systems such as our pressure activated binding organic liquids (BOLs) showed that gas uptake had a marked dependence on the polarity of the sorbent itself.8 This polarity dependence is needed because of the equilibrium between the nonpolar (gas lean, nonionic) and polar (gas rich, ionic) species (Figure 1). As a rule of thumb, acid-gas loading increases with the polarity of the solvent. We tested four commercially available tertiary dialkanolamines: methyldiethanolamine (MDEA), ethyldiethanolamine (EDEA), diisopropylethanolamine (DIPEA), and di-n-butylethanolamine (DBEA). Of the alkanolamines, MDEA gelled and became extremely viscous after CO2 uptake, but while DIPEA and DBEA were less viscous they also showed poor gas uptake due to a lower polarity. Thus, EDEA was chosen because of its gravimetric capacity and lower viscosity increase compared to other tertiary dialkanolamines. Phase equilibrium data were needed to demonstrate the effectiveness of CO2 and H2S absorption by anhydrous tertiary alkanolamines under gas processing conditions. For this reason, we set out to study the potential selectivity of H2S over CO2 under high-pressure gas mixtures by first studying their behavior with gaseous species commonly found in natural gas. Binary vapor−liquid−liquid equilibrium (VLLE) measurements were performed for CH4, C3H8, H2S, and CO2 with anhydrous EDEA at 10−50 °C by the TPx and TPxy methods. We present here the first published VLLE measurements of CH4, C3H8, H2S, and CO2 in anhydrous EDEA and assess its phase equilibria and trends with each gas. We also present preliminary process simulations using a molecular-thermodynamic model developed in Aspen Plus to evaluate anhydrous EDEA’s performance compared to other solvents such as Fluor Solvent, Selexol, and MDEA in a typical high-CO2 natural gas absorber.
Figure 2. Schematic of glass TPx and TPxy apparatus.
measured using a McLeod gauge or a mercury barometer. The pressures were measured with an accuracy of ±0.07 kPa. The temperature of the mixture in the cell was measured using a platinum resistance thermometer inserted into a thermowell in the top of the cell, and was measured with an accuracy of ±0.05 °C. The TPx data at pressures above 100 kPa were measured in the apparatus shown schematically in Figure 3. The apparatus consisted of a 300 cm3 stainless steel cylinder immersed in a constant temperature oil bath. Lines were connected to the cell
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PHASE EQUILIBRIUM DATA The apparatuses and techniques used for the TPx and TPxy measurements are similar to those used by Giles et al.9 A brief description of the TPx and TPxy apparatuses and procedures follows. The reported TPx data at pressures below about 100 kPa for the EDEA + H2S system were measured using the apparatus shown schematically in Figure 2. This apparatus consisted of a 300 cm3 glass cell and manometer in a constant temperature bath. The cell was sealed with a PTFE cap and O-ring. The cell was equipped with charging and degassing lines. The pressure was measured using a mercury manometer attached to the cell. The pressure of the reference side of the manometer was
Figure 3. Schematic of stainless steel TPx apparatus. 17563
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same glass apparatus as the low-pressure TPx measurements (Figure 2). A set of measurements TPxy was initiated by charging the desired amounts of EDEA and H2S or CO2 to the cell. The cell was placed in the constant temperature bath, and the bath was brought to the desired run temperature. The cell contents were stirred for 30 min or longer to allow the phases to come to equilibrium. At 50 °C, the viscosity was low enough that we had little difficulty in stirring these mixtures. At 10 °C, the viscosity was very high, making it difficult to stir. It also appeared that the viscosity increased as the H2S or CO2 was added to the EDEA. In the CO2 system, the stirring became so difficult that it was decided to attempt the measurements at 23.1 °C where the viscosity would be lower. Stirring improved, but was still difficult, and the system required at least an hour for the pressure to equilibrate. The stirrer was then shut off and multiple samples of the liquid phase were analyzed. For the CO2 systems, the samples were analyzed using a flash gas/flash liquid technique. The sample was withdrawn into a weighed test tube attached to a buret filled with water. As the sample was removed, the dissolved CO2 “flashed” and displaced the water in the buret. The volume of CO2 was then measured. The test tube was then weighed to determine the amount of flash liquid in the test tube. With corrections for the residual CO2 dissolved in the flash liquid and the vapor pressure of EDEA, the overall composition of the liquid phase was determined. In a few cases, the vapor phase was sampled and analyzed for EDEA, and the analyses came back as “nondetect.” This further justifies the conclusion of this work that the vaporphase concentration of EDEA is negligibly small in all the measurements. For the H2S measurements, the liquid samples were analyzed using an Anton Paar density meter. Mixtures of EDEA and H2S were gravimetrically made up at different compositions. The densities of these mixtures were measured at 20 °C and correlated (using an excess volume model). The results of these measurements are shown in Table 1. The densities of the samples from the TPxy cell were then measured at 20 °C. Using the density correlation, the composition of a sample could then be determined.
for adding material, degassing, and measuring the pressure. The temperature was measured using a platinum resistance thermometer to within ±0.05 °C. The pressure was measured using a Paroscientific pressure transducer with an uncertainty of ±0.7 kPa. A TPx run was initiated by evacuating the cell and then charging it with a known amount of EDEA. The cell was then submerged in an isothermal bath and agitated until equilibrium was established at the desired temperature. The cell contents were then degassed by withdrawing vapor from the cell into a weighed, evacuated receiver. Degassing was done to remove any compounds from the cell that were lighter than the compounds of interest. These light compounds can cause significant errors in the total pressure measurements. Increments of the second compound were then added to the cell. After each addition, the cell contents were brought to equilibrium at the desired temperature. The pressure was then measured and recorded. This procedure was repeated until over half of the entire composition range had been traversed. The other half of the composition range was covered by repeating the same procedure with different amounts of EDEA and solute. Where applicable, the vapor pressures of the solutes were measured as a separate charge. The reported TPx measurements above 100 kPa were measured using the apparatus shown schematically in Figure 4.
Table 1. H2S + EDEA Density Measurements at 20 °C; Used in TPxy Analysis Figure 4. Schematic of visual TPxy apparatus.
The visual cell consisted of a piece of glass pipe (50 mm i.d.) with stainless steel flanges on each end. Lines for pressure measurement and for sampling the liquid and vapor phases extended through the top of the cell. A calibrated platinum resistance thermometer was placed into a thermowell which extended into the cell for temperature measurement. The liquid in the cell was stirred with a Teflon coated magnetic stir bar. The cell had an internal volume of approximately 300 cm3. The system pressure was measured with a Paroscientific digital pressure transducer with a calibrated uncertainty of ±0.7 kPa. The platinum thermometers were calibrated using the ice and steam points of water and were referenced to an NIST traceable standard resistance thermometer. The temperatures are estimated to have an accuracy of ±0.05 °C. The reported TPxy measurements for EDEA + H2S at 10 °C were performed in the
wt fraction (H2S)
density, g/mL
0 0.0129 0.041 0.052 0.081
1.0140 1.0202 1.0334 1.0379 1.0501
Tables 2 and 3 show the results of the TPx measurements for methane and propane at 10 and 50 °C. These tables show the measured pressure in kilopascals, the molar charge composition (z), and the calculated molar liquid composition (x). The measured pressure is assumed to be the solute partial pressure since the vapor pressure of EDEA is negligibly small. Tables 4 and 6 report the results of the TPx and TPxy measurements for the CO2 + EDEA system at the three temperatures studied. The tables for the TPx measurements show the measured pressure in kilopascals, the molar charge composition (z), and the molar liquid composition (x). The 17564
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Table 2. TPx Data for the CH4−EDEA Binary System at 10 and 50 °Ca
a
Table 4. TPx Data for the CO2−EDEA Binary System at 10 and 50 °Ca
temp, °C
press, kPa
z(CH4), mole fraction
x(CH4), mole fraction
temp, °C
press, kPa
z(CO2), mole fraction
x(CO2), mole fraction
10 10 10 10 10 10 10 10 50 50 50 50 50 50 50 50
9317.5 7958.2 6403.1 4908.0 3309.8 1677.1 1130.7 567.5 12518.7 10429.1 8290.6 6209.8 4138.0 2068.0 1385.8 695.5
0.1550 0.1366 0.1159 0.0891 0.0625 0.0323 0.0213 0.0102 0.1550 0.1366 0.1159 0.0891 0.0625 0.0323 0.0213 0.0102
0.0722 0.0645 0.0556 0.0435 0.0310 0.0162 0.0108 0.0052 0.0967 0.0847 0.0716 0.0547 0.0382 0.0197 0.0130 0.0062
10 10 10 10 10 10 10 10 10 10 10 10 10 10 50 50 50 50 50 50 50 50 50 50
4503.9 4499.4 4486.2 4497.8 4491.5 4234.3 2959.8 3028.1 1977.8 1978.6 1398.8 584.9 82.71 33.74 6217.1 5566.3 3997.5 3575.2 2621.7 2436.5 1777.7 1473.5 695.4 332.5
1 0.9052 0.8053 0.5501 0.4931 0.7065 0.3948 0.5964 0.5046 0.3186 0.4274 0.2255 0.1290 0.0695 0.4932 0.7065 0.3948 0.5964 0.3187 0.5046 0.4274 0.2255 0.1290 0.0695
1 0.8984 0.7066 0.4811 0.3866 0.3522 0.3182 0.3174 0.2840 0.2709 0.2547 0.2039 0.1235 0.0681 0.5055 0.4625 0.3559 0.3350 0.2727 0.2630 0.2134 0.1862 0.1050 0.0567
The vapor phase is assumed to be pure CH4.
Table 3. TPx Data for the C3H8−EDEA Binary System at 10 and 50 °Ca temp, °C
press, kPa
z(C3H8), mole fraction
x(C3H8), mole fraction
10 10 10 10 10 10 10 10 10 10 10 10 10 10 50 50 50 50 50 50 50 50 50 50 50 50 50 50
638.4 636.9 637.3 636.3 637.9 636.5 637.7 636.5 637.7 636.3 638.3 637.5 386.1 206.6 1700.6 1697.3 1713.1 1711.3 1696.8 1695.5 1710.7 1700.2 1709.9 1509.9 1206.5 1295.9 724.3 365.6
1 0.8997 0.7972 0.6965 0.5553 0.5964 0.5004 0.5084 0.4056 0.4398 0.3012 0.2046 0.0952 0.0402 1 0.8997 0.7972 0.6965 0.5553 0.5004 0.5964 0.4056 0.3012 0.5084 0.4398 0.2046 0.0952 0.0402
1 0.8988 0.7876 0.6700 0.5526 0.5438 0.4954 0.4244 0.3954 0.3251 0.2838 0.1790 0.0749 0.0282 1 0.8978 0.7684 0.6066 0.5507 0.4891 0.3955 0.3791 0.2525 0.1905 0.1576 0.1560 0.0616 0.0220
a
a
The vapor phase is assumed to be pure CO2.
Table 5. TPx Data for the H2S−EDEA Binary System at 10 and 50 °Ca
The vapor phase is assumed to be pure C8H8.
measured pressure is assumed to be the CO2 partial pressure since the vapor pressure of EDEA is negligibly small. The tables for the TPxy measurements show the measured pressure in kilopascals and the measured liquid composition (x). As with the TPx measurements, the vapor is assumed to be all CO2. In these tables, the reported compositions represent the
a
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temp, °C
press, kPa
z(H2S), mole fraction
x(H2S), mole fraction
10 10 10 10 10 10 10 10 10 10 10 10 10 50 50 50 50 50 50 50 50 50 50 50 50 50
1371.3 1377.4 1367.0 1106.5 823.8 538.9 525.0 366.4 320.0 162.9 28.88 15.22 9.79 3545.7 3488.4 3366.1 2112.7 1707.8 1342.9 1218.1 801.3 673.4 591.6 391.7 207.5 85.5
1 0.9151 0.8324 0.7589 0.7016 0.6406 0.6171 0.5954 0.5626 0.4646 0.2319 0.1265 0.0728 1 0.9099 0.8067 0.7072 0.6119 0.6193 0.5540 0.5293 0.4496 0.4690 0.3323 0.2104 0.0958
1 0.9142 0.8242 0.7429 0.6826 0.6217 0.6134 0.5781 0.5590 0.4620 0.2259 0.1223 0.0699 1 0.9089 0.7817 0.6645 0.6050 0.5747 0.5461 0.4887 0.4417 0.4321 0.3252 0.2045 0.0925
The vapor phase is assumed to be pure H2S.
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Table 6. TPxy Data for the CO2−EDEA Binary System at 10, 23.1, and 50 °Ca
a
temp, °C
press, kPa
x(CO2), mole fraction
10 10 23.1 23.1 50 50 50
1056.0 688.4 217.0 157.3 1021.0 736.8 366.2
0.2609 0.2471 0.1601 0.1333 0.1543 0.1155 0.0688
complexes with effectively zero vapor pressure, and the complex is assumed to have zero charge.
overall or apparent composition before accounting for any complexation, i.e., the apparent compositions. Tables 5 and 7 report the results of the TPx and TPxy measurements for H2S + EDEA at the two temperatures studied. Table 7. TPxy Data for the H2S−EDEA Binary System at 10 and 50 °Ca
a
press, kPa
x(H2S), mole fraction
10 10 10 50 50 50
34.27 24.51 16.15 202.8 158.9 106.7
0.2479 0.2191 0.1536 0.2203 0.1774 0.1219
(1)
H 2S + EDEA ↔ H 2S ·EDEA
(2)
Equation 1 assumes that 1 mol of CO2 complexes or solvates with 2 mol of EDEA, while the H2S solvation is assumed to follow 1:1 stoichiometry. The stoichiometries in eqs 1 and 2 were guided empirically by best fits of the data. Evaluation of spectroscopic data6,7 suggests that complexation of both CO2 and H2S follow 1:1 stoichiometry, but the CO2 loading is too low to confirm the 1:1 stoichiometry. We have accepted the empirically determined 1:2 stoichiometry for CO2 as phenomenologically acceptable. An activity-coefficient model has been used to describe the physical nonideality of the mixture, and in particular, the NRTLRK property option in Aspen Plus has been chosen in which the NRTL15 model is used for the activity coefficients and the Redlich−Kwong16 (RK) equation of state was used to estimate the vapor-phase departures from the ideal-gas law. Further details on the NRTL model are available in the textbook by Prausnitz et al.17 All species, including supercritical methane, have been treated in the symmetrical convention. The model includes the Poynting correction to evaluate the effect of pressure on the pure-component fugacities. The molar volumes, needed for the Poynting correction, for CH4, C3H8, CO2, H2S, and EDEA were taken to be temperature-independent as 0.090, 0.076, 0.037, 0.036, and 0.132 m3/kmol, respectively. The molar volumes for all components except CH4 are equal to the value at the respective boiling points; for CH4, the molar volume is relatively high and was determined to fit the data presented in Table 2. The overall thermodynamic modeling framework adopted here may be considered to be the “solvated-solution-with-physicalinteractions” framework presented in the textbook by Prausnitz et al.17 Pure-component parameters have been used unchanged from Aspen Plus, version 7.3. EDEA properties are uncertain, but they are not important for the present application. The vapor pressure of EDEA is negligibly small, and its mole fraction in the vapor phase is effectively zero. Later developments of the technology will reevaluate this assumption in case solvent losses need to be estimated. The measured vapor−liquid equilibrium (VLE) data presented in Tables 2−7 have been used to fit optimum values of the NRTL and chemical-equilibrium parameters by using the Aspen Plus Data Regression System. The NRTL parameters are presented in Table 8. The nomenclature here follows that in the Aspen Plus documentation. The NRTL model has two binary parameters: the nonrandomess parameter (αij), which is assumed to be symmetric and independent of temperature, and the energy parameter (τij), which is asymmetric and is assumed to vary linearly with inverse temperature. For binaries not presented in Table 8, the energy parameters have been defaulted to zero.
The vapor phase is assumed to be pure CO2.
temp, °C
CO2 + 2EDEA ↔ CO2 · EDEA 2
The vapor phase is assumed to be pure H2S.
These tables show the same information as is contained in Tables 4 and 6 for the CO2 system. We note that the measured pressure is assumed to be the H2S partial pressure since the vapor pressure of EDEA is negligibly small. For the TPx measurements, the reported charge compositions were determined from the amount of each component weighed into the cell for that point. The reported liquid compositions are obtained using a constant volume flash calculation where the vapor phase is assumed to be 100% solute with no EDEA. In this flash calculation, the liquid density was calculated using the Rackett equation with the reference density reported in Table 1. The liquid densities of propane, CO2, and H2S were taken from the correlations reported in the DIPPR database.10 The bulk liquid volume was calculated assuming no volume of mixing. This assumption introduced negligible error in determining the liquid composition.
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THERMODYNAMIC MODELING Thermodynamic correlations for chemical absorption are typically phenomenological or semiempirical models,11−13 but they may also be totally empirical.14 Here, we use a semiempirical approach in which the chemical absorption of CO2 and that of H2S by EDEA are modeled by assuming that they form solvated
Table 8. Fitted NRTL Parameters for the CH4−C3H8−CO2−H2S−EDEA System, Including Complexes component 1 component 2 a12 a21 b12 b21 α12 = α21
CO2 EDEA 0 0 1143 −216 0.250
CO2 CO2·EDEA2, H2S·EDEA 0 0 1280 1280 0.250
H2S EDEA 0 0 2 3819 0.250
H2S CO2·EDEA2, H2S·EDEA 0 0 2051 −617 0.250 17566
CH4 EDEA, CO2·EDEA2, H2S·EDEA 0 0 496 496 0.200
C3H8 EDEA, CO2·EDEA2, H2S·EDEA 0.75 −1.83 1035 996 0.351
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bij (3)
T
The logarithms of the equilibrium constants for the two reactions (eqs 1 and 2) are assumed to vary linearly with inverse temperature. ln(K1) ≡
ln(K 2) ≡
[CO2 ·EDEA 2] 2
[CO2 ][EDEA]
= −15.42 +
4892 T
[H 2S ·EDEA] 3457 = −9.89 + [H 2S][EDEA] T
(4)
(5)
In eqs 4 and 5, the square brackets denote component activity, which is equal to the product of the mole fraction and the activity coefficient. It is noted that the number of parameters in the two equilibrium constants (eqs 4 and 5) and NRTL parameters (Table 8) have been carefully minimized to avoid overfitting the data. Only linear dependence in 1/T has been used for the logarithm of the two equilibrium constants. Temperature dependence of the energy binary parameter has only been allowed for the EDEA−C3H8 system. As far as possible, the energy parameters have been forced to be symmetric. Finally, the NRTL parameters for CH4 + EDEA, CH4 + CO2·EDEA2, and CH4 + H2S·EDEA have been forced to be the same, and the same constraint has been applied to the NRTL parameters for C3H8 with EDEA and the two complexes; see the last two columns of Table 8. The restricted parameter set has not limited the accuracy of the thermodynamic model and is expected to facilitate reliable extrapolation. Extended parameters may be introduced in the future as the database increases. The enthalpy of solution has been estimated from the phaseequilibrium model using the Gibbs−Helmholtz equation.18 ⎛ ∂H ⎞ ⎛ ∂ ln fi ⎞ − hi* ≈ R ⎜ ΔH̅ 0i ≡ ⎜ ⎟ ⎟ ⎝ ∂(1/T ) ⎠σ ,{x 0} ⎝ ∂n0i ⎠T , P , n 0j
Figure 5. TPx diagram in the methane−EDEA system. Comparison of experimental data (points) to model calculations (lines).
measured data at 10 °C and the squares represent the data at 50 °C. The NRTL-RK model (lines) fit provides an excellent representation of the experimental data. It should be noted that this phase behavior is for acid-gas-lean EDEA, which is less polar than EDEA solvated with CO2 or H2S. Methane is estimated be less soluble in either the produced EDEA alkylcarbonate or hydrosulfide ionic liquids. The model calculations for the methane−EDEA binary are judged to be adequate for process analysis. Propane−EDEA. Propane uptake by EDEA is low, but, as for CH4, not as low as with aqueous amines. The measured and calculated solubilities of propane by anhydrous EDEA at 10 and 50 °C are shown in Figure 6. At pressures above 600 kPa at 10 °C
(6)
In eq 6, ΔH̅ 0i is the differential enthalpy of solution of component i, which is defined as the difference between the partial molar enthalpy of component i and its ideal-gas enthalpy at the same temperature, hi* is the ideal-gas enthalpy of component i at the same temperature, R is the gas constant, f i is the fugacity of component i, and the derivative on the right side of eq 6 is taken along the equilibrium line at constant apparent loading. Mathias and O’Connell18 have shown that eq 6 is approximate, but is expected to be accurate for the liquid phase at conditions far from the critical point, and this approximation is valid for the present application. The pure-component parameters of the two complexes (in particular, the enthalpies of formation and the heat capacity) have been adjusted such that the enthalpies of solution of CO2 and H2S estimated by eq 6 are accurately represented by the Aspen Plus model.
Figure 6. TPx Diagram in the propane−EDEA system. Comparison of experimental data (points) to model calculations (lines).
and 1700 kPa at 50 °C, propane−EDEA begins to exhibit liquid− liquid equilibrium (LLE). Similar to the methane case, this VLLE plot is presented for the acid-gas-lean EDEA, which is less polar than EDEA reacted with CO2 or H2S. VLLE data have not been collected for propane with EDEA containing either CO2 or H2S, but the uptake of propane in such systems is expected to be lower due to the increased polarity of the produced alkylcarbonate or hydrosulfide ionic liquids. The agreement between model calculations and data is good. The concentration of EDEA in liquid propane is small, but the data do not help determine the quantitative solubility. The NRTL parameters have been adjusted such that the solubility of
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ANALYSIS OF BINARY PHASE BEHAVIOR In this section the binary phase behavior data presented in Tables 2−7 are analyzed, and the ability of the thermodynamic model to accurately correlate the data and provide insight into the underlying phenomena is studied. Methane−EDEA. Methane uptake by EDEA is low, but not as low as for aqueous amines. Figure 5 shows a plot of total pressure versus composition for the CH4 system over pure EDEA at 10 and 50 °C. In Figure 5, the diamonds represent the 17567
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too low, but it is noted that the TPxy data are accurately represented, albeit at lower uptakes. The model accurately represents the LLE at 10 °C; however, the predicted LLE pressure is biased somewhat low. The model predictions of the LLE at 23 °C do not have experimental data for comparison. Figure 9 presents model estimations of the fraction of CO2 that is solvated. The solvation fraction decreases with temperature
EDEA in liquid propane is less than 1 mol %, but the accuracy of this estimate is not known at the present time. CO2−EDEA. Figure 7 presents experimental data and model calculations for the absorption of CO2 in anhydrous EDEA.
Figure 7. TPx diagram for the CO2−EDEA system: log−log scale. The points show measured TPx and TPxy data at 10, 23, and 50 °C. The lines present model calculations at the three temperatures.
Figure 9. Fraction of CO2 that is solvated, as estimated by the model. The solid circles represent the maximum liquid mole fraction, beyond which LLE occurs, as estimated by the data and model.
There is fairly good agreement between the TPx and TPxy data at 50 °C, but at lower temperatures, there are discrepancies between the TPx and TPxy measurements. The main cause for the discrepancies is the high solution viscosity at the lower temperatures. The model provides an overall best fit of the TPx and TPxy data sets. The experimental and calculated results indicate that the system exhibits liquid−liquid equilibrium (LLE) at high pressures and when CO2 is subcritical (Tc = 31 °C). At 10 °C, LLE occurs when the apparent mole fraction of CO2 is about 0.38 and the pressure reaches about 4400 kPa. Figure 8 has the same information content as Figure 7, but is on a linear scale to better reveal the liquid−liquid equilibrium. Figure 8 demonstrates the complex phase behavior exhibited by the CO2−EDEA system, but also highlights the problems with accurate measurement and modeling. The calculations of CO2 uptake at 50 °C seem to be
and drops when the CO2 mole fraction exceeds about 1/3, which corresponds to a CO2 loading of 0.5. Figure 9 identifies the maximum mole fraction at 10 and 23 °C, beyond which LLE occurs. The model predictions in Figure 9 are consistent with experimental observations.19 Figure 10 presents the enthalpy of
Figure 10. Estimated enthalpy of solution of CO2 in EDEA.
solution estimated by the model. The calculated enthalpies of solution are consistent with the extent of solvation results shown in Figure 9, because we expect a more negative enthalpy of solution when the extent of solvation is higher. The magnitude of the enthalpy of solution decreases with increasing temperature, and drops off sharply as the extent of solvation decreases (corresponding to a CO2 mole fraction of about 1/3). As a comparison, the enthalpy of solution of CO2 in aqueous MEA is about −85 kJ/mol18 under absorber conditions (∼40 °C and CO2 loading of
