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Solubility of Carbon Dioxide in Aqueous Solutions of Monoethanolamine in the Low and High Gas Loading Regions Michael Wagner, Inga von Harbou, Jaeik Kim, Irina Ermatchkova, Gerd Maurer, and Hans Hasse* Laboratory of Engineering Thermodynamics, University of Kaiserslautern, P.O. Box 30 49, D-67653 Kaiserslautern, Germany S Supporting Information *

ABSTRACT: Reliable data for the solubility of carbon dioxide in aqueous solutions of monoethanolamine are required for the design and evaluation of postcombustion carbon capture processes. As published experimental data for the solubility of carbon dioxide in aqueous solutions of monoethanolamine show considerable scatter, the solubility of carbon dioxide in aqueous solutions containing (15 and 30) mass percent of monoethanolamine; that is, (2.9 and 7.0) mol·(kg H2O)−1 respectively, was measured at molar ratios of carbon dioxide to monoethanolamine in the liquid solution from 0.1 to 1.3 at (313, 353 and 393) K. An apparatus based on headspace gas chromatography (on the synthetic gas solubility method) was used for the experiments at low (high) gas loadings, that is, at partial pressures of carbon dioxide from (1 to 80) kPa (from (0.4 to 8.6) MPa). The new experimental results are compared to literature data and used to parametrize a physicochemical thermodynamic model based on the extended Pitzer equation for the Gibbs excess energy of the liquid mixture. Furthermore, model predictions for the ion speciation in the liquid phase are compared to literature data from NMR spectroscopy.

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INTRODUCTION The postcombustion capture (PCC) of carbon dioxide (and its storage) is one of the most important methods in the reduction of CO2 emissions from coal and natural gas fired power plants. Currently absorption of CO2 by aqueous solutions of amines is the most advanced PCC technology.1,2 An aqueous solution of monoethanolamine (MEA)in particular a solution with about 30 mass percent of MEAis considered to be the benchmark solvent for this application. The design of the reactive absorption process, especially of the absorber and the desorber, requires reliable data for thermodynamic properties, especially for the solubility of CO2 in such solutions. Many publications report experimental data for the solubility of CO2 in aqueous solutions of MEA (cf. Table 1). The results from various groups, however, scatter considerably especially in the case of the 30 mass percent solution of MEA in water. The differences in the experimental results for the partial pressure of CO2 above the same solution reach up to 1 order of magnitude. Therefore, currently, it is nearly impossible to develop and parametrize a reliable thermodynamic model for the solubility of CO2 in aqueous solutions of MEA. Continuing previous work on the solubility of sour gases in aqueous solutions of Nmethyldiethanolamine (MDEA) and/or piperazine (in the low pressure region by headspace gas chromatography;3−5 in the high pressure region by the synthetic gas solubility method6−11), this paper reports new experimental data for the solubility of CO2 in aqueous 15 and 30 mass percent solutions of MEA in water at (313, 353, and 393) K at low and high gas loadings (stoichiometric molar ratios of CO2 to MEA © 2013 American Chemical Society

between 0.1 and 1.3) and consequently low and high partial pressures of CO2 (between (0.001 and 8.6) MPa).

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EXPERIMENTAL SECTION Experimental Techniques. The partial pressure of CO2 in a postcombustion carbon capture process that uses an aqueous solution of MEA typically varies between about 0.4 and 15 kPa in the absorber12 and reaches up to 200 kPa in the desorber.13 Experimental investigations at higher pressures, for example, at pressures up to 10 MPa are nevertheless of great interest as previous investigations on the solubility of sour gases in aqueous solutions of MDEA and/or piperazine showed that a thermodynamic model for such gas solubility phenomena should be primarily based on experimental work in the high pressure range.6−11 As carbon dioxide is predominantly dissolved chemically (i.e., as bicarbonate or as carbamate) while the amine is converted to ionic species, interactions between the reaction products are important in the liquid phase. Experimental results in the high pressure region allow a more reliable determination of the corresponding interaction parameters as there the concentration of the reaction products is high. Therefore, a sound physicochemical thermodynamic model for the gas solubility based on such experimental data also gives good predictions for the low partial pressure region. However, experimental data in the low pressure region can be Received: September 20, 2012 Accepted: February 28, 2013 Published: April 2, 2013 883

dx.doi.org/10.1021/je301030z | J. Chem. Eng. Data 2013, 58, 883−895

Journal of Chemical & Engineering Data

Article

interactions between neutral (unreacted) amine and the reaction products). As there is no single experimental technique available to reliably measure the partial pressure of CO2 over aqueous solutions of MEA from 1 kPa to 10 MPa, two different experimental techniques were used in the present investigation: A synthetic gas solubility technique was used for investigations at total pressures between (0.4 and 8.7) MPa and headspace gas chromatography was applied for investigations at CO2 partial pressures between (1 and 80) kPa. Synthetic Gas Solubility Apparatus. The synthetic gas solubility apparatus enables a reliable determination of the pressure that is required to dissolve a precisely known amount of CO2 at a fixed temperature in a precisely known amount of solvent. The experimental setup and the experimental procedures have been described in detail in previous publications.11,14,15 Hence the main features are restated here. The differences resulting from the particular system of the present investigation are described. Figure 1 shows a scheme of the equipment.

Table 1. Literature Review of CO2 Solubility Data in Aqueous Solutions of MEA T/K

molality author Mason and Dodge53 1936 Reed and Wood54 1941 Lyudkovskaya and Leibush55 1949 Atadan56 1954 Muhlbauer and Monaghan57 1957 Jones et al.17 1959 Goldman and Leibush27 1959 Murzin and Leites58 1971 Lee et al.20 1974 Lawson and Garst24 1976 Lee et al.18 1976 Nasir and Mather59 1977 Isaacs et al.60 1980 Chan H.M.61 1981 Maddox et al.19 1987 Austgen et al.16 1991 Shen and Li22 1992 Dawodu and Meisen62 1994 Jou et al.30 1995 Song et al.23 1996 Jane and Li25 1997 Park et al.21 1997 Mathonat et al.63 1998 Park et al.64 2002 Bonenfant et al.65 2003 Dang and Rochelle66 2003 Ma’mun et al.31 2005 Hilliard51 2005 Harris et al.67 2009 Portugal et al.32 2009 Puxty et al.68 2009 Kadiwala et al.26 2010 Zhao et al.69 2010 Arcis et al.70 2011 Aronu et al.28 2011

mol·(kg H2O)

pCO2/kPa

min

max

0.1−54.3

273

348

0.2

100

2.9

373

413

150

1700

0.5; 2.3; 7.2

298

348

255

4124

0.5; −25.7 2.8−3.0

303 298

343 373

112 0.1

2620 145

3.0 1.0; 2.2; 3.0; 7.2

313 348

413 413

0.1 0.8

931 470

0.5 - 4.3

303

353

0.003

300

3.0; 7.2 2.9

313 313

373 413

0.1 1.3

6616 2787

1.1; 3; 4.9; 7.2 2.5; 5.0

298 353

393 373

0.1 0.01

10 3.0

3.0 0.1 3.0

353 298 298

373 298 353

0.01 0.1 5.6

2 0.1 6792

3.0

313

353

0.1

229

3.0 5.6

313 373

313 373

15.7 455

2550 3863

7.0 3.0 3.0 3.0 7.0

273 313 353 313 313

423 313 353 313 393

0.1 3.1 3.6 3.5 2000a

20000 2359 122 2092 10000a

1.8; 3 0.9

313 296

313 296

2.6 100a

2189 100a

2.9; 7.2

313

333

0.02

9.6

7.0

393

393

7

192

3.5−11.2 7.0 2.9; 3

313 313 313

333 313 313

0.006 90a 0.2

50 2510a 63

1.1 7.0

313 313

313 313

3.2 116

908 5329

7.0 2.9; 7.0 2.9; 7.0; 13.4; 24.6 2.9

303 323 313

303 373 393

1500a 530 0.001

1500a 5190 1060

313

313

1

27

373 313 313

443 393 393

12 0.032 4

1626 411 408

Mazinani et al.71 2011b Xu et al.72 2011 6.9−7.0 Brúder et al.29 2012 7.2 Tong et al.73 2012 7.0 a

−1

min

max

Figure 1. Synthetic gas solubility apparatus for measuring the solubility of a single gas in a solvent at elevated pressures: A, cylindrical high-pressure equilibrium view cell with two sapphire windows and magnetic stirrer; B, thermostat; C, gas container; D, pressure transducers; E, tank for rinsing water; F, tank for solvent mixture; G, high-pressure spindle press; H, AC-bridge with three platinum resistance thermometers; I, solution outlet; J, cooling trap; K, vacuum pump.

The central part is a horizontally placed, cylindrical high pressure view cell (volume of about 30 cm3) with sapphire windows on both ends. It is thermostatted by silicone oil flowing through an annular jacket. The jacket of the view cell is insulated to reduce heat loss to the surroundings. At the beginning of a measurement the previously evacuated cell is charged with carbon dioxide from a small gas container. In all experiments, the amount of CO2 in the view cell (between 1.6 g and 6.5 g) is precisely determined by weighing the gas container before and after the charging process with a high precision balance. In some experiments at 393 K, the amount of CO2 in the view cell was also determined volumetrically (from temperature, pressure, and the volume of the cell). Then the solvent was added via a high-pressure spindle press which was

Total pressure. bNo numerical data available.

used to test the extrapolation capability of the model and to additionally determine some model parameters that are not accessible from high pressure gas solubility data (such as for 884



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filled with the aqueous solution of MEA. Typically, about 1.2 dm3 of solution were prepared gravimetrically in an evacuated glass storage tank from two glass storage tanks that were filled with degassed MEA and water, respectively. The amounts of water and MEA were determined gravimetrically. The volume displacement of the piston of the spindle press is measured and used (together with the solvent density that is measured by a vibrating tube densimeter (Anton Paar GmbH, model DMA 4500 M)) to determine the amount of added solvent. That amount is always somewhat above the minimum amount that is required to completely dissolve the gas; that is, after equilibration the view cell is filled with a homogeneous liquid phase at a pressure above the solubility pressure. Then very small amounts of the liquid mixture are withdrawn (by the high-pressure spindle press) until the first tiny stable gas bubbles appear. The solubility pressure is approximated by the pressure when the first tiny stable bubbles are observed. The pressure was measured with two high precision pressure transducers (WIKA GmbH, Klingenberg, Germany; full scale (2.5 and 10) MPa, respectively; intrinsic uncertainty: 0.1 % of the transducer’s full scale) in connection with a mercury barometer (Lambrecht, Göttingen, Germany). The transducers were calibrated against a high-precision pressure balance (Desgranges & Huot, Aubervilliers, France) before and after each measurement series. The uncertainty of the experimental results (for the solubility pressure) is the sum of the intrinsic uncertainty of the pressure transducer and a correction for the experimental procedure (expansion of the cell volume from a homogeneous system to a vapor−liquid equilibrium). Numbers are given with the experimental data. The solvent is an aqueous solution of MEA. Its composition is known from its gravimetric preparation with an absolute uncertainty of ± 0.0024 mol MEA·(kg H2O)−1. The relative uncertainty of the mass of solvent in the view cell is about 0.13 %. The relative uncertainty of the stoichiometric molality of CO2 (i.e., the amount of CO2 per kilogram of water) in the liquid phase is between 0.34 % and 0.67 %. Two calibrated platinum resistance thermometers in the thermostatted jacket of the view cell were used to determine the temperature with an estimated uncertainty of less than ± 0.1 K. Headspace Gas Chromatography. Headspace gas chromatography was applied to determine the solubility of CO2 in the low gas loading region. The experimental setup was described in previous publications.3−5 Therefore, only the main features are repeated here. Figure 2 shows a scheme of the headspace gas chromatography arrangement. The main components are eight thermostatted sample cells (only one sample cell is shown in Figure 2), a vapor phase expansion and sampling system and a gas chromatograph. In an experiment, the sample cells (stainless steel vials, volume, 30 cm3) are filled with about 25 cm3 of the loaded solvent. The sample cells are mounted in a cell holder where they are equilibrated for at least 12 h. After equilibration, the sample cell is pressurized with nitrogen from buffer tank A to a constant pressure ((0.2, 0.26, and 0.45) MPa for the measurements at (313, 353, and 393) K, respectively). Then, the vapor phase is expanded to a larger volume by connecting the vapor phase of the cell to buffer tank B which is pressurized to a lower constant pressure ((0.17, 0.2, and 0.37) MPa for the measurements at (313, 353, and 393) K, respectively). The sample loop (SL; internal volume, 20 μL) is filled in that step. The sample valve is then switched and the sample is transferred to the gas chromatograph. After each measurement, the sampling system

Figure 2. Scheme of the headspace gas chromatography arrangement: CH, liquid-thermostatted cell holder (temperature T1); VH, liquidthermostatted valve holder (temperature T2 > T1); A, nitrogen tank (higher pressure); B, nitrogen tank (lower pressure); GC, gas chromatograph; He, helium (carrier gas); SC, sample cell; MV, multiposition valve; S1 to S8, sample positions; P1 to P8, purge positions; SV, sample valve; SL, sample loop.

is purged with nitrogen. A multiposition valve (Valco Instruments Co. Inc., type 2CSD16MWE) allows a connection to each of the eight sample cells (by stainless steel capillaries, inner diameter, 0.5 mm) to the sample loop (positions S1 to S8). The other eight positions (P1 to P8) are used for purging. The multiposition valve and the sample valve are mounted in a sample valve holder. Both valves are operated pneumatically by an electronic controller. The temperature of the valve holder is kept at 15 K above the temperature of the cell holder, and the line to the gas chromatograph is also kept at a higher temperature to avoid condensation. The gas chromatograph (Agilent, type 6890) is equipped with a capillary column (Alltech, type Heliflex AT-Q, 30 m, 0.32 mm i.d.) and a thermal conductivity detector. The peak area resulting from carbon dioxide in the sample is used to determine the partial pressure of CO2 above the liquid solution in the sample cell. For calibration, a sample cell is filled with pure CO2 (at a pressure ranging from about 5 kPa to 80 kPa). The pressure is measured with a high precision pressure transducer (MKS, Andover, Massachusetts, US, type 690A13TRA) with an accuracy of ± 0.05 %. The maximum relative deviation between the resulting calibration line and a single calibration point was below 2 %. In a gas chromatographic experiment between 4 and 8 sample cells were filled with the same loaded liquid mixture. The standard deviation between the peak area of CO2 from the different cells was typically below 2 % but under some unfavorable conditions, it reached 7 %. The composition of the liquid mixture in the sample cell is known from its preparation. Typically, about 1.2 dm3 solution were prepared gravimetrically in an evacuated glass storage tank by dissolving known amounts of MEA in water. About 250 cm3 of that unloaded solution were transferred to another evacuated storage tank (volume ≈ 300 cm3) where it was charged with CO2. The masses of both the unloaded solution and the solute CO2 were determined gravimetrically with an accuracy of ± 2 mg. The storage tanks were then shaken for about (4 to 5) h and stored at room temperature for about 24 h before the sample cells of the headspace equipment were charged with the loaded solutions. The stoichiometric molalities of MEA and CO2 in water were calculated from the masses of the single 885



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components applying some corrections for the presence of vapor phases in the tanks. These corrections are generally small (< 0.03 % for MEA and < 0.6 % for CO2). The uncertainty of each of these corrections is estimated to ± 20 %. The total relative uncertainty of the amount of CO2 in the liquid phase ranges from about ± 0.04 % at 313 K to about ± 0.6 % at 393 K. It was estimated from the uncertainty of the gravimetric sample preparation (calculated with a Gauss error propagation) and the estimated uncertainty of the correction that was applied to account for the vapor phases. The total relative uncertainty in the molality of MEA in the liquid phase in a sample cell is estimated to be smaller than ± 0.1 %. The temperature is measured with a calibrated platinum resistance thermometer in the liquid that is used for thermostatting the sample cell holder. The estimated uncertainty of the temperature measurement is ± 0.1 K. Materials. Carbon dioxide (4.5, volume fraction ≥ 99.995 %) was purchased from Messer Griesheim GmbH, Krefeld, Germany. MEA was purchased from Sigma-Aldrich (mass fraction ≥ 99 %). For all evaluations the impurities in MEA were neglected. Ultrapure water was taken from a water purification system (TWF/EI-Ion UV Plus TM, Siemens AG, Water Technologies). The liquid components were degassed before they were used for the preparation of the solvent mixtures. Experimental Results and Comparison with Literature Data. The solubility of CO2 in aqueous solutions of MEA was measured at T ≈ 313 K, 353 K, and 393 K. Table 2 gives an

Table 3. Solubility of Carbon Dioxide in Aqueous Solution of MEA (m̅ MEA ≈ 2.9 mol·(kg H2O)−1) in the High Pressure Range (ΔT = ± 0.1 K, Δm̅ MEA/m̅ MEA = ± 0.1 %) T

Table 2. Investigated Temperatures T, Stoichiometric Molalities m̅ , Loadings αCO2 and Partial Pressures of CO2 pCO2

T K 313 353 393

αCO2

pCO2

molCO2/molMEA

kPa

m̅ MEA mol·(kg H2O)−1

min

max

min

max

2.9 7.0 2.9 7.0 2.9 7.0

0.45 0.47 0.21 0.29 0.07 0.10

1.28 1.04 1.11 0.87 0.95 0.72

2.3 1.4 1.6 1.6 2.1 2.9

8070 7510 6310 7360 8550 7060

K


m̅ CO2 mol·(kg H2O)−1

MPa

ptot

313.63 313.27 313.22 313.21 313.11 313.20 353.14 353.23 353.15 353.18 353.19 353.20 353.24 353.20 353.17 353.23 393.11 393.16 393.15 393.14 393.15 393.17 393.15 393.15 393.16 393.16

2.856 2.911 2.911 2.911 2.889 2.931 2.930 2.930 2.930 2.930 2.930 2.930 2.930 2.930 2.930 2.930 2.856 2.856 2.856 2.856 2.856 2.856 2.856 2.856 2.856 2.856

2.863 ± 0.019 2.962 ± 0.019 3.125 ± 0.019 3.216 ± 0.019 3.541 ± 0.02 3.760 ± 0.02 2.111 ± 0.019 2.156 ± 0.019 2.244 ± 0.019 2.374 ± 0.019 2.548 ± 0.02 2.665 ± 0.02 2.732 ± 0.02 2.836 ± 0.02 2.899 ± 0.02 3.080 ± 0.02 1.518 ± 0.019 1.720 ± 0.019 1.894 ± 0.019 1.992 ± 0.019 2.087 ± 0.019 2.17 ± 0.019 2.308 ± 0.02 2.425 ± 0.02 2.52 ± 0.02 2.699 ± 0.02

1.813 ± 0.014 2.134 ± 0.011 2.77 ± 0.02 3.45 ± 0.07 5.24 ± 0.02 8.07 ± 0.02 0.803 ± 0.011 0.970 ± 0.011 1.198 ± 0.011 1.623 ± 0.011 2.40 ± 0.02 3.10 ± 0.02 3.47 ± 0.02 4.34 ± 0.03 4.75 ± 0.02 6.35 ± 0.03 0.786 ± 0.011 1.359 ± 0.011 2.111 ± 0.011 2.49 ± 0.02 3.30 ± 0.02 3.65 ± 0.02 4.67 ± 0.02 5.44 ± 0.02 6.58 ± 0.02 8.74 ± 0.02

H2O)−1). Most of the results available in the literature are given as loading versus partial pressure of CO2. The loading is given as m̅ CO2 αCO2 = m̅ MEA (1) For the comparisons shown in Figures 3 and 4, the partial pressure of CO2 was calculated from the experimental results for the total pressure by subtracting the partial pressures of water and MEA, which were calculated (for given temperature and composition of the liquid phase) from the model described below. The new experimental results for the partial pressure of CO2 above aqueous solutions of MEA agree with literature data within the (wide) scattering of the published data. As a lot of literature data is available, the detailed comparison is restricted to a few typical examples. The model described below (with temperature specific interaction parameters) was used for interpolating the results of the present work to the conditions (i.e., the composition of the liquid phase) of the literature data. As shown in Figure 3 some of the literature data for the partial pressure of CO2 above a 15 mass percent aqueous solution of MEA agree well with the new experimental results, whereas some literature data lie systematically above and some other data systematically below the new experimental data. For example, the agreement is good with the data reported by Austgen et al.,16 Jones et al.,17 Lee et al.,18 and Maddox et al.19 where the deviations are below 10 %. Examples for experimental data which lie systematically above the new

overview over the experiments. Two solvent mixtures were investigated. The mass fraction (molality) of MEA in the solvent was 0.15 (2.9 mol·(kg H2O)−1) and 0.30 (7.0 mol·(kg H2O)−1). The loading of the aqueous solution, that is, the molar ratio of CO2 to MEA, ranges from 0.1 to 1.3 and the partial pressure of CO2 from 1 kPa to 8550 kPa. The experimental results from the investigation with the synthetic gas solubility technique (i.e., for the total pressure ptot that is required to dissolve carbon dioxide in the aqueous solution of MEA) are listed in Table 3 (for the solvent with m̅ MEA ≈ 2.9 mol·(kg H2O)−1) and Table 4 (for the solvent with m̅ MEA ≈ 7.0 mol·(kg H2O)−1) together with their estimated experimental uncertainties. The experimental results from the investigations with the synthetic method are shown and compared to literature data in Figure 3 (for the solvent with m̅ MEA ≈ 2.9 mol·(kg H2O)−1) and Figure 4 (for the solvent with m̅ MEA ≈ 7.0 mol·(kg 886



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Table 4. Solubility of Carbon Dioxide in Aqueous Solution of MEA (m̅ MEA ≈ 7.0 mol·(kg H2O)−1) in the High Pressure Range (ΔT = ± 0.1 K, Δm̅ MEA/m̅ MEA = ± 0.1 %) K

T


313.23 313.24 313.24 313.18 313.23 313.26 313.24 313.23 313.25 313.20 353.26 353.16 353.17 353.18 353.16 353.15 353.21 353.18 393.14 393.19 393.19 393.19 393.21 393.18 393.19 393.21 393.19 393.16

7.008 7.008 7.008 7.069 7.008 7.008 7.008 7.008 7.008 7.008 7.069 7.069 7.069 7.069 7.069 7.069 7.069 7.069 7.069 7.069 7.069 7.069 7.069 7.069 7.069 7.069 7.069 7.069


MPa

± ± ± ± ± ± ± ± ± ± ± ± ± ± ± ± ± ± ± ± ± ± ± ± ± ± ± ±

0.448 ± 0.030 0.695 ± 0.011 0.739 ± 0.011 0.745 ± 0.011 1.312 ± 0.011 1.671 ± 0.011 3.62 ± 0.02 4.92 ± 0.02 7.32 ± 0.02 7.51 ± 0.05 0.513 ± 0.011 0.981 ± 0.011 1.715 ± 0.011 2.234 ± 0.015 2.71 ± 0.02 3.80 ± 0.02 5.52 ± 0.04 7.40 ± 0.02 0.644 ± 0.011 0.972 ± 0.022 1.249 ± 0.011 2.084 ± 0.031 2.76 ± 0.02 3.23 ± 0.02 3.41 ± 0.02 4.21 ± 0.02 5.23 ± 0.02 7.23 ± 0.03

4.980 5.267 5.261 5.405 5.739 5.965 6.627 6.855 7.233 7.279 4.151 4.544 4.862 5.027 5.182 5.457 5.732 6.059 3.370 3.603 3.731 3.995 4.148 4.255 4.316 4.483 4.709 4.987

0.023 0.023 0.023 0.023 0.024 0.024 0.024 0.024 0.025 0.024 0.023 0.024 0.023 0.024 0.024 0.024 0.024 0.025 0.024 0.024 0.024 0.025 0.022 0.022 0.022 0.022 0.023 0.023

ptot

results are the data reported by Lee et al.20 (deviations of about 20 %), by Park et al.21 (deviations of about 16 %), by Shen and Li22 (deviations of about 15 %), and by Song et al.23 (deviations of about 17 %). Examples for experimental data which lie systematically below the new results are the data reported by Lawson and Garst24 (deviations of about 18 %) and by Jane and Li25 (deviations of about 14 %). As shown in Figure 4, literature data for the partial pressure of CO2 above aqueous solutions with m̅ MEA ≈ 7.0 mol·(kg H2O)−1, reveal also a large scattering. At 313 K, the experimental data from the present investigation agree best with the data published by Kadiwala et al.26 (deviations of about 8 %). At 353 K the best agreement is observed with the data by Lee et al.18 (deviations of about 7 %). Most literature data reported for 393 K were determined at low gas loadings (αCO2 < 0.5), whereas the new experimental results are for higher gas loadings (0.5 < αCO2 < 0.7). The data reported by Goldman and Leibush,27 Aronu et al.,28 Brúder et al.29 for αCO2 < 0.49 (pCO2 < 0.45 MPa) agree well with the new experimental resultsthe relative deviations are below 7 %. But the few literature data reported for 0.5 < αCO2 < 0.7 reveal deviations of up to 40 % from the new experimental results for the partial pressure of CO2. The experimental results from the investigations with the headspace gas chromatography (i.e., at low loadings (partial pressures of CO2)) are listed in Table 5 (m̅ MEA ≈ 2.9 mol·(kg H2O)−1) and Table 6 (m̅ MEA ≈ 7.0 mol·(kg H2O)−1). The

Figure 3. Partial pressure of carbon dioxide above liquid mixtures of (CO2 + MEA + H2O), m̅ MEA = 2.9 mol·(kg H2O)−1 at T ≈ 313 K (top), 353 K (middle), 393 K (bottom). Experimental results: ■, this work, synthetic method; ×, Austgen et al.;16 ◑, Jane and Li;25 ○, Jones et al.;17 ⊕, Lawson and Garst;24 □, Lee et al.;20 Δ, Lee et al.;18 , Maddox et al.;19 ◓, Park et al.;21 ☆, Park et al.;64 ◊, Shen and Li;22 +, Song et al.23

experimental results are reported together with their estimated uncertainties. The uncertainty of the experimental results for the partial pressure of CO2 was estimated from the uncertainty of the corresponding peak area and an additional contribution from the conversion of the peak area to the partial pressure of CO2 via the calibration curve. This uncertainty was estimated to 2 % or 0.4 kPa, which ever is larger. The last column of Table 5 and Table 6 reports the standard deviation ΔpCO2,repr of the single results for pCO2 (as obtained when 4 or 8 sample cells were filled with the same mixture). As expected, for most experiments, ΔpCO2,repr is less than the estimated experimental uncertainty ΔpCO2. This supports the quality of the estimation of the experimental uncertainties. The experimental results from the investigations with headspace gas chromatography are shown and compared to literature data in Figure 5 (m̅ MEA ≈ 2.9 mol·(kg H2O)−1) and 887



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Table 5. Solubility of Carbon Dioxide in Aqueous Solution of MEA (m̅ MEA = 2.9 mol·(kg H2O)−1) in the Low Pressure Range (ΔT = ± 0.1 K, Δm̅ MEA/m̅ MEA = ± 0.1 %) K

T


313.16 313.16 313.18 313.16 313.17 313.16 313.17 313.18 353.16 353.17 353.17 353.17 353.17 353.17 353.17 353.17 353.16 393.20 393.16 393.17 393.14 393.18 393.11 393.19 393.17 393.18 393.17

2.863 2.863 2.869 2.863 2.869 2.863 2.869 2.869 2.872 2.893 2.893 2.872 2.893 2.872 2.893 2.872 2.893 2.818 2.818 2.999 2.999 2.999 2.999 2.818 2.999 2.818 2.999


pCO2

ΔpCO2,repr

kPa

kPa

± ± ± ± ± ± ± ± ± ± ± ± ± ± ± ± ± ± ± ± ± ± ± ± ± ± ±

2.30 ± 0.4 5.94 ± 0.4 9.93 ± 0.4 14.7 ± 0.4 21.7 ± 0.4 28.1 ± 0.6 39.9 ± 0.8 56.2 ± 1.1 1.59 ± 0.4 2.79 ± 0.4 5.17 ± 0.4 8.01 ± 0.4 12.8 ± 0.4 20.4 ± 0.4 23.9 ± 0.5 49.4 ± 1.0 59.9 ± 1.2 2.06 ± 0.4 3.49 ± 0.4 5.65 ± 0.4 11.3 ± 0.4 19.9 ± 0.4 26.0 ± 0.5 49.1 ± 1.0 38.9 ± 0.8 74.0 ± 1.5 58.2 ± 1.2

0.02 0.08 0.15 0.3 0.5 0.3 0.8 0.2 0.02 0.06 0.10 0.10 0.2 0.1 0.1 0.3 0.2 0.02 0.04 0.07 0.3 0.2 0.3 0.4 0.3 0.9 0.4

1.417 1.521 1.585 1.638 1.689 1.732 1.788 1.840 0.750 0.872 1.029 1.120 1.226 1.322 1.364 1.492 1.531 0.205 0.315 0.439 0.612 0.753 0.829 0.932 0.934 1.020 1.038

0.001 0.002 0.002 0.002 0.002 0.003 0.003 0.003 0.001 0.001 0.001 0.001 0.001 0.002 0.002 0.002 0.002 0.001 0.001 0.001 0.002 0.001 0.001 0.002 0.002 0.002 0.002

K from Lee et al.18 (typical relative deviation for the partial pressure of CO2: at 353 K (393 K) 32 % (17 %)), and a good agreement with the experimental results by Jones et al.17 (7 % (10 %)). At 313 K the new experimental results for the partial pressure of CO2 above a 7 m (30 mass percent) aqueous solution of MEA agree best with the data reported by Aronu et al.28 typical relative deviation, 18 %. At 353 K, the new experimental results lie between the data reported by Jou et al.30 (which are lower by about 24 %) and those reported by Aronu et al.28 (which are higher by about 22 %). At 393 K the scattering of the literature data is comparably small and the new experimental results agree well with the data by Goldman and Leibush27 (typical deviation 17 %), Lee et al.18 (typical deviation 25 %), Ma’mun et al.31 (typical deviation 20 %) and Brúder et al.29 (typical deviation 18 %). The density of unloaded aqueous solutions of MEA is required for the evaluation of the gas solubility measurements by the synthetic method. The density of both aqueous solutions was determined with a vibrating tube densimeter at temperatures between (293 and 320) K and correlated. The correlation results are given in Table 7. Primary data are given in the Supporting Information document.

Figure 4. Partial pressure of carbon dioxide above liquid mixtures of (CO2 + MEA + H2O), m̅ MEA = 7.0 mol·(kg H2O)−1 at T ≈ 313 K (top), 353 K (middle), 393 K (bottom). Experimental results: ■, this work - synthetic method; ◒, Aronu et al.;28 ◐, Brúder et al.;29 ☆, Goldman A.M.;27 ⬡, Harris et al.;67 ▽, Jou et al.;30 , Kadiwala et al.;26 □, Lee et al.;20 Δ, Lee et al.;18 ×, Mathonat et al.;63 ◊, Shen and Li;22 ◑, Tong et al.;73 ⊕, Xu et al.72

Figure 6 (m̅ MEA ≈ 7.0 mol·(kg H2O)−1). As in the high pressure range, also in the low pressure range the new data lie within the wide scatter of the literature data. Most methods used in literature18,20,22,30,31 determined the CO2-concentration in the liquid phase by precipitation with BaCl2 and subsequent titration. This method, however, has an estimated error of at least 3 %30 and is thus less reliable than the gravimetric determination of the CO2 concentration used in the present work. The detailed comparison with literature data is restricted to a few examples (cf. the discussion of Figures 3 and 4). At 313 K, the new experimental results for the partial pressure of CO2 above a nearly 3 m aqueous solution of MEA agree well with the experimental results by Austgen et al.16 (relative deviations in the partial pressure of CO2: 3 %), Park et al.21 (8 %) and Shen and Li22 (14 %) and fairly with the results by Aronu et al.28 (24 %) and Portugal et al.32 (19 %). There is a fair agreement with the experimental results at 353 K and 393

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THERMODYNAMIC MODEL The experimental data were correlated using a thermodynamic model. The model used to describe the electrolyte solution is similar to the model used previously in our group.3−8,11,33,34 A 888



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Table 6. Solubility of Carbon Dioxide in Aqueous Solution of MEA (m̅ MEA ≈ 7.0 mol·(kg H2O)−1) in the Low Pressure Range (ΔT = ± 0.1 K, Δm̅ MEA/m̅ MEA = ± 0.1 %) K

T


313.20 313.16 313.21 313.21 313.21 313.21 313.16 313.21 352.86 352.83 352.81 352.83 352.85 352.81 352.84 352.82 352.83 352.83 352.83 391.99 391.99 391.98 392.03 392.00 391.94 392.00 391.98 392.00 392.01 392.01 391.96

7.055 7.097 6.919 7.055 7.097 7.097 7.035 6.919 7.192 6.992 7.036 6.992 7.192 7.036 6.992 6.992 6.992 7.019 7.019 7.109 7.338 7.338 7.251 7.338 7.109 7.338 7.338 7.338 7.109 7.251 7.251


pCO2

ΔpCO2,repr

kPa

kPa

± ± ± ± ± ± ± ± ± ± ± ± ± ± ± ± ± ± ± ± ± ± ± ± ± ± ± ± ± ± ±

1.38 ± 0.4 3.27 ± 0.4 6.08 ± 0.4 9.02 ± 0.4 17.5 ± 0.4 29.2 ± 0.6 49.1 ± 1.0 66.4 ± 1.3 1.58 ± 0.4 3.75 ± 0.4 6.59 ± 0.4 11.8 ± 0.4 18.4 ± 0.4 23.0 ± 0.5 29.6 ± 0.6 39.7 ± 0.8 55.2 ± 1.1 70.4 ± 1.4 78.7 ± 1.6 2.92 ± 0.4 3.10 ± 0.4 4.24 ± 0.4 7.16 ± 0.4 12.2 ± 0.4 16.0 ± 0.4 24.2 ± 0.5 27.6 ± 0.6 41.6 ± 0.8 62.4 ± 1.3 69.8 ± 1.4 73.2 ± 1.5

0.03 0.03 0.05 0.17 0.3 0.5 1.4 0.8 0.03 0.20 0.35 0.3 0.2 0.2 0.6 0.4 1.8 0.7 0.5 0.04 0.10 0.17 0.10 0.3 0.6 0.6 2.0 2.1 2.4 1.2 0.3

3.296 3.498 3.519 3.679 3.826 3.958 4.061 4.106 2.105 2.462 2.787 3.002 3.204 3.228 3.289 3.400 3.489 3.591 3.622 0.730 0.845 0.969 1.239 1.564 1.704 2.004 2.087 2.369 2.497 2.618 2.633

0.001 0.001 0.002 0.002 0.002 0.003 0.004 0.005 0.002 0.002 0.001 0.002 0.002 0.002 0.002 0.002 0.003 0.003 0.003 0.002 0.002 0.002 0.002 0.002 0.002 0.002 0.002 0.002 0.003 0.003 0.003

Figure 5. Partial pressure of carbon dioxide above liquid mixtures of (CO2 + MEA + H2O), m̅ MEA = 2.9 mol·(kg H2O)−1 at T ≈ 313 K (top), 353 K (middle), 393 K (bottom). Experimental results: ●, this work - headspace gas chromatography; ◒, Aronu et al.;28 ×, Austgen et al.;16 ◐, Isaacs et al.;60 ◑, Jane and Li;25 ○, Jones et al.;17 ⊕, Lawson and Garst;24 □, Lee et al.;20 Δ, Lee et al.;18 , Maddox et al.;19◓, Park et al.;21 ☆, Park et al.;64 ⊞, Portugal et al.;32 ◊, Shen and Li;22 +, Song et al.23

crude outline of the model is given below. More details are presented in the Supporting Information document. Vapor−Liquid Equilibrium. The vapor−liquid equilibrium condition is formulated using the extended Henry’s law for i = CO2 and MEA on the molality scale: ⎡ νi∞(p − p s ) ⎤ (m) W ⎢ ⎥ai(m) = pyi φi exp kH, i RT ⎣ ⎦

(2)

(3)

chemical potentials of all solute species are normalized according to Henry’s law on the molality scale, the activity of a solute species i is m ai(m) = i γi(m) (4) m°

In eqs 2 and 3 and aW are the vapor phase mole fraction and the liquid phase activities of component i (i = CO2 or MEA) and water, respectively. k(m) H,i is Henry’s constant of volatile solute i in pure water (based on the molality scale) at the vapor pressure of pure water psW. ν∞ i and vW are the partial molar volume of component i infinitely diluted in water and the molar volume of pure liquid water, respectively. R is the universal gas constant. The influence of pressure on ν∞ i and νW is neglected. As the molality scale is used to express the composition of the liquid phase (with water being the solvent) and the

where γ(m) is the activity coefficient (on the molality scale) of i solute species i that is calculated from an expression for the excess Gibbs energy GE of the liquid mixture. The activity of water is calculated from the expression for GE using the Gibbs− Duhem equation (cf. below). The reference state for the chemical potential of water is the pure liquid at temperature and pressure of the system, whereas for a solute species it is a hypothetical 1 molal solution of that solute in pure water, mi = m° = 1 mol·(kg H2O)−1, at the temperature and pressure of the system where it experiences the same interactions as in infinite dilution in pure water.

and the extended Raoult’s law for water: ⎡ ν W (p − p s ) ⎤ s W ⎥aW = pyW φW φWs exp⎢ pW RT ⎣ ⎦

yi, a(m) i

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there are several reversible chemical reactions in aqueous solutions of MEA and CO2, the speciation of the liquid phase has an essential influence on the vapor−liquid equilibrium. Therefore, the chemical reaction equilibrium has to be taken into account. Chemical Reaction Equilibria in Aqueous Solution of CO2 and MEA. The following (reversible) chemical reactions are accounted for the liquid phase: The autoprotolysis of water (H2O):

H 2O ⇌ H+ + OH−

(I)

The formation of bicarbonate (HCO−3 ): CO2 + H 2O ⇌ HCO−3 + H+

(II)

The formation of carbonate (CO2− 3 ): HCO−3 ⇌ CO32 − + H+

(III)

The deprotonation of monoethanolamine (MEA): MEAH+ ⇌ MEA + H+

(IV)

The formation of carbamate (MEACOO−): MEA + HCO−3 ⇌ MEACOO− + H 2O

(V)

Table 8 gives the chemical structures of MEA and its reaction products. The liquid phase is a mixture of CO2, H2O, and MEA and all ionic species occurring in chemical reactions I to V, whereas the vapor phase is treated as a mixture of CO2, H2O, and MEA. The chemical equilibrium condition for a chemical reaction r (= I to V) is given by eq 5.

K r (T ) =

∏ aiν

i ,r

i

where Kr is the chemical reaction equilibrium constant for reaction r and νi,r is the stoichiometric factor of reactant i in reaction r (νi,r > 0 for a product and νi,r < 0 for a reactant). Kr is assumed to depend only on temperature; that is, the influence of pressure is neglected. Water is treated as the solvent while CO2 and MEA and all ionic species are treated as solutes, that is, ai → a(m) i . The thermodynamic equilibrium is solved considering the boundary conditions (mass balance and chemical reaction equilibrium) for each stoichiometric component i (i = CO2, H2O, and MEA) to calculate the speciation of the liquid phase (i.e., the true molalities mi of all solute species i) for given temperature T and stoichiometric molalities of MEA (m̅ MEA) and CO2 (m̅ CO2) in the liquid phase. Gibbs Excess Energy Model. Pitzer’s molality-scale based equation for the excess Gibbs energy of aqueous electrolyte solutions is used to calculate the activity coefficients of all solute species as well as the activity of water.35−37 Details are given in the Supporting Information document. Required Thermodynamic Properties. Chemical Reaction Equilibrium Constants. The temperature-dependent chemical reaction equilibrium constants Kr(T) (r = I to V) were obtained from literature.38−43 They are given in Table 9. The influence of temperature on the equilibrium constant for the protonation of MEA (reaction IV) was determined from the experimental data of Bates and Pinching41 (273 K to 323 K) and Bénézeth et al.42 (273 K to 573 K). The formation of carbamate (reaction V) was investigated by McCann,43 Böttinger,44 and McCann et al.45 The correlation equation

Figure 6. Partial pressure of carbon dioxide above liquid mixtures of (CO2 + MEA + H2O), m̅ MEA = 7.0 mol·(kg H2O)−1 at T ≈ 313 K (top), 353 K (middle), 393 K (bottom). Experimental results: ●, this work - headspace gas chromatography; ◒, Aronu et al.;28 ◐, Brúder et al.;29 +, Dang and Rochelle;66 ○, Hilliard;51 ☆, Goldman A.M.;27 ▽, Jou et al.;30 □, Lee et al.;20 Δ, Lee et al.;18 ⊞, Ma’mun et al.;31 ◊, Shen and Li;22 ◑, Tong et al.;73 ⊕, Xu and Rochelle.72

Table 7. Density ρ of Aqueous Solution of MEA; ρ/kg·dm−3 = c + d·T/K ; T = (293 to 320) K; Δρ = ± 0.0004 kg·dm−3 m̅ MEA mol·(kg H2O)−1

c

d

2.89 7.01

1.1194 1.1544

−3.9037 × 10−4 −4.8057 × 10−4

(5)

Table 8. Chemical Structures of Reactants of Reactions IV and V

φsW and φi are the fugacity coefficients of pure water at saturation and the fugacity coefficients of component i (i = CO2, H2O, or MEA) in the vapor mixture, respectively. As 890



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Table 9. Chemical Reaction Equilibrium Constants (on the Molality Scale): ln Kr = A + B/(T/K) + C· ln(T/K) + D·(T/K) + E/ (T/K)2

a

reaction

A

B

C

D

E

I II III IV V

140.932 −1203.01 175.36 −1.73782 −5.9680

−13445.9 68359.6 −7230.6 −6092.85 2888.6

−22.4773 188.444 −30.6509 0 0

0 −0.206424 0.01315 0.001157 0

0 −4712910 −372805 0 0

T/K 298 273 273 273 293

to to to to to

source Fisher and Barnes38 Patterson et al.39 Patterson et al.40 this worka McCann43

573 673 523 398 333

Based on experimental data by Bates and Pinching41 (273 K to 323 K) and Bénézeth et al.42 (273 K to 573 K).

Table 10. Henry’s Constants of Component i in Water (on the Molality Scale): ln(k(m) H,i /MPa) = A/(T/K) + B·ln(T/K) + C·(T/ K) + D component i

A

B

C

D

T/K

source

CO2 MEA

−9624.4 −11803.5

−28.749 −10.617

0.01441 0

192.876 84.599

273 to 473 288 to 408

Rumpf and Maurer15 this work

Table 11. Partial Molar Volume of CO2 in Water at Infinite 46 Dilution ν∞ CO2

Table 14. Summary of Deviations (dev) between Experimental Data and Correlation Resultsa synthetic method

T

ν∞ CO2

K

cm3/mol

313 353 393

33.4 36.3 40.8

T/K

mean rel dev p

max abs dev pCO2/kPa

mean rel dev pCO2

15 ma %

313 353 393

4.1 % 2.1 % 5.1 %

0.9 1.9 0.7

7.0 % 7.7 % 4.7 %

30 ma %

313 353 393

2.8 % 6.9 % 7.0 %

0.5 0.7 0.4

6.4 % 9.3 % 1.9 %

Table 12. Second Virial Coefficients49,50 T

BCO2,CO2

BCO2,H2O

BH2O,H2O

K

cm3/mol

cm3/mol

cm3/mol

313 353 393

−112 −85.1 −64.7

−163 −129 −104

−930 −578 −385

mean overall

β(0) CO2,MEACOO− β(0) MEA,MEAH+ β(0) MEA,MEACOO− β(0) HCO3−,MEAH+ β(0) MEAH+,MEACOO− mCO2,HCO3−,MEAH+

Apar −6.998·10−2 −6.982·10−2

Bpar 30.55 42.08

2.9872·10−1

−49.92

−2.361·10−2 2.2865·10−1 4.352·10−2

50.94 −104.45 −14.39

1.9653·10−1 −2.131·10−2

6.2 %

For the total pressure p (synthetic method) and the partial pressure of CO2 pCO2 (head space GC). For partial pressures below 10 kPa the maximal absolute deviation in pCO2 is reported. The other numbers are mean relative deviations. The last line gives the mean value of the corresponding deviations over all measurements.

Table 13. Interaction Parameters in Pitzer’s G Equation. Par = Apar + Bpar/(T/K) for T/K = 313 to 393 parameter

4.7 %

a

E

β(0) MEA,MEA β(0) CO2,MEAH+

head space GC

The partial molar volume of CO2 dissolved in water at infinite dilution (ν∞ CO2) was calculated as recommended by Brelvi and O’ Connell.46 The numerical values are given in Table 11. Owing to the lack of experimental data the partial molar volume of MEA at infinite dilution in water (ν∞ MEA) was neglected; that is, it was set to zero. Properties of Pure Water. The vapor pressure of pure water psW was calculated from the equation by Saul and Wagner.47 The relative permittivity εW, which is required in Pitzer’s equation for the excess Gibbs energy, was adopted from Bradley and Pitzer.48 Vapor Phase Fugacity Coefficients. The virial equation of state, which was truncated after the second virial coefficient, was used to calculate all vapor phase fugacity coefficients. Purecomponent second virial coefficients were calculated from a correlation based on data recommended by Dymond.49 All second virial coefficients were neglected when (at least) one of the interacting species is MEA. The second virial coefficient for the mixture of CO2 and H2O was calculated as recommended by Hayden and O’ Connell.50 All non-neglected second virial coefficients are given in Table 12.

−44.05 5.888

for reaction V is based on experimental data for T ≤ 333 K. It was also used at higher temperatures. Henry’s Constant for the Solubility of CO2 and MEA in Water and Partial Molar Volumes at Infinite Dilution of CO2 and MEA in Water. The correlation equation for Henry’s constant of CO2 in water was taken from Rumpf and Maurer.15 Henry’s constant of MEA in water was determined following the procedure described by Ermatchkov and Maurer34 (for Nmethyldiethanolamine (MDEA) instead of MEA). Details are given in the Supporting Information document. The correlation equations for Henry’s constants k(m) H,i of CO2 and MEA (on the molality scale and at the vapor pressure of pure water) are given in Table 10. 891



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Figure 7. Total pressure (top) and partial pressure of carbon dioxide (bottom) above liquid mixtures of (CO2 + MEA + H2O). Left, m̅ MEA = 2.89 mol·(kg H2O)−1; right, m̅ MEA = 7 mol·(kg H2O)−1. Experimental results, headspace gas chromatography: ▽, 313 K; ◊, 353 K; Δ, 393 K. Experimental results, synthetic method: ▼, 313 K; ⧫, 353 K; ▲, 393 K. Correlation: −.

Interaction Parameters in Pitzer’s GE Equation. As there are many solutes species in an aqueous solution of (MEA + CO2), there is a large number of binary and ternary interaction parameters in the extended Pitzer equation for the excess Gibbs energy. The number of parameters is too large to be all adjusted meaningfully to the experimental data. However, only a few interaction parameters are really important. These parameters were selected in preliminary studies before the final adjustment of parameters was done. A detailed parameter study revealed that the most important interaction parameters depend on temperature. Therefore, in a first correlation step, the experimental results for each temperature were correlated to achieve a proper method for interpolations in the concentration at constant temperature. However, for the design of separation equipment, a correlation over a wide temperature range is required. As the new experimental results were determined only for three temperatures, only two-parameter equations are reasonable to describe the influence of temperature on these parameters. The influence of temperature on an interaction parameter Par was described by Par = A par + Bpar /(T /K )

Tables 5 and 6)) is 6.2 %. A detailed summary is given in Table 14. A detailed study revealed that for some interaction parameters eq 6 is not well suited to describe the new experimental data within experimental uncertainty. When the parameters were fit to the experimental data for the single isotherm the average relative deviation decreased by up to 30 %. Therefore, further experimental investigations at other temperatures are recommended to achieve a better correlation for the influence of temperature on the interaction parameters. A comparison between the correlation results and the new experimental results is shown in Figure 7. Speciation StudyComparison with NMR Data. The speciation of the liquid mixture has an essential influence on the vapor−liquid equilibrium (and related properties). The chemical reactions in the liquid phase determine the speciation. Calculation results for the speciation can be compared with experimental results from investigations by NMR-spectroscopy by Böttinger,44 Hilliard,51 and Jakobsen et al.52 Figure 8 shows some comparisons between the experimental results of those groups for a 7 m aqueous solution of MEA at (293, 313, 333, and 353) K and calculation results. The investigations by Hilliard51 were restricted to low gas loadings (αCO2 < 0.5), whereas Böttinger44 and Jakobsen et al.52 reported speciation data also for higher gas loadings (αCO2 < 1.1). The experimental uncertainty of the NMR-spectroscopic investigations is difficult to assess. An indication is given by the difference of the overall amine concentration in the solution as determined from the sample preparation and the overall amine

(6)

The resulting parameters are given in Table 13. The average relative deviation between the new experimental results for the total pressure above aqueous solutions of (MEA + CO2) from the synthetic method (see Tables 3 and 4) and the correlation results is 4.7 % and for the partial pressure of CO2 (at pCO2 > 10 kPa from the headspace chromatographic investigation (see 892



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Figure 8. Speciation in the mixture of (CO2 + MEA + H2O), m̅ MEA = 7 mol·(kg H2O)−1 at T ≈ 293 K, 313 K, 333 K, and 353 K (left to right and top to bottom). Experimental results: ●, Böttinger;44 ○, Jakobsen et al.;52 □, Hilliard.51 Prediction: −, this work.

■

concentration determined from the NMR spectra (by summing up the contributions of the different species). Jakobsen et al.52 report relative uncertainties of up to 24 %. The corresponding number for the measurements of Böttinger44 exceeds 10 % in particular at higher gas loadings. Figure 8 shows some typical comparisons between calculation results and NMR-spectroscopic data for the speciation in a 7 m aqueous solution of MEA at temperatures from 293 K to 353 K. The experimental results of Bö t tinger 44 are shown together with experimental uncertainties that were estimated from the differences between the NMR-results and the gravimetric preparation of the mixtures. For the lower temperatures most calculation results agree with Böttinger’s experimental data within the estimated uncertainty. The speciation data from experimental investigations by different authors reveal similar deviations as those between the calculation results of the present work and the experimental data by Bö ttinger.44 The above-mentioned authors also extended their investigations to MEA molalities between (2.9 and 4.1) mol·(kg H2O)−1. Also in that concentration range, the difference between experimental and calculation results are of the same order of magnitude as mentioned above. In the Supporting Information document a comparison between the measured and calculated CO2 species (except MEACOO−) is shown (cf. Figure S.1).

CONCLUSION

Two different experimental techniques were applied to measure the solubility of CO2 in two aqueous solutions of MEAthe mass fraction of MEA in the unloaded aqueous solutions was (15 and 30) mass percentat three temperatures (313, 353 and 393) K. Headspace gas chromatography was applied to measure the partial pressure of CO2 above the loaded aqueous solutions in the low gas loading range, whereas the synthetic gas solubility technique was applied to determine the total pressure above such solutions in the high gas loading range. The new experimental data agree with literature data within the (wide) scattering of the literature data. A previously developed thermodynamic framework was adapted to describe the new experimental data. The model provides a sound representation at the three temperatures investigated. Model predictions for the speciation are also in good agreement with (the limited) literature data from NMR-spectroscopic investigations. However, more experimental investigations of the thermodynamic properties of such systems (for example, on the solubility of CO2 in aqueous solutions of MEA at different temperatures as well as for different compositions of the binary (MEA + water) system), and on the chemical reaction equilibria, as well as calorimetric and spectroscopic investigations might result in further improvements of the model. 893



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(15) Rumpf, B.; Maurer, G. An experimental and theoretical investigation on the solubility of carbon dioxide in aqueous solutions of strong electrolytes. Ber. Bunsen. Phys. Chem. 1993, 97, 85−97. (16) Austgen, D. M.; Rochelle, G. T.; Chen, C. C. Model of vapor− liquid equilibria for aqueous acid gas-alkanolamine systems. 2. Representation of hydrogen sulfide and carbon dioxide solubility in aqueous MDEA and carbon dioxide solubility in aqueous mixtures of MDEA with MEA or DEA. Ind. Eng. Chem. Res. 1991, 30, 543−555. (17) Jones, J. H.; Froning, H. R.; Claytor, E. E. Solubility of acidic gases in aqueous monoethanolamine. J. Chem. Eng. Data 1959, 4, 85− 92. (18) Lee, J. I.; Otto, F. D.; Mather, A. E. Equilibrium between carbon-dioxide and aqueous monoethanolamine solutions. J. Appl. Chem. Biotechnol. 1976, 26, 541−549. (19) Maddox, R.; Bhairi, A.; Diers, J.; Thomas, P. Equilibrium solubility of carbon dioxide or hydrogen sulfide in aqueous solutions of monoethanolamine, diglycolamine, diethanolamine, and methyldiethanolamine. GPA Res. Rep. 1987, 104, 1−47. (20) 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, 803−805. (21) Park, S.-B.; Shim, C.-S.; Lee, H.; Lee, K.-H. Solubilities of carbon dioxide in the aqueous potassium carbonate and potassium carbonatepoly(ethylene glycol) solutions. Fluid Phase Equilib. 1997, 134, 141−149. (22) Shen, K.-P.; Li, M.-H. Solubility of carbon dioxide in aqueous mixtures of monoethanolamine with methyldiethanolamine. J. Chem. Eng. Data 1992, 37, 96−100. (23) Song, J.-H.; Yoon, J.-H.; Lee, H.; Lee, K.-H. Solubility of carbon dioxide in monoethanolamine + ethylene glycol + water and monoethanolamine + poly(ethylene glycol) + water. J. Chem. Eng. Data 1996, 41, 497−499. (24) Lawson, J. D.; Garst, A. W. Gas sweetening data: Equilibrium solubility of hydrogen sulfide and carbon dioxide in aqueous monoethanolamine and aqueous diethanolamine solutions. J. Chem. Eng. Data 1976, 21, 20−30. (25) Jane, I.-S.; Li, M.-H. Solubilities of mixtures of carbon dioxide and hydrogen sulfide in water + diethanolamine + 2-amino-2-methyl1-propanol. J. Chem. Eng. Data 1997, 42, 98−105. (26) Kadiwala, S.; Rayer, A. V.; Henni, A. High pressure solubility of carbon dioxide (CO2) in aqueous piperazine solutions. Fluid Phase Equilib. 2010, 292, 20−28. (27) Goldman, A. M.; Leibush, A. G. Study of the equilibrium of carbon dioxide desorption from monoethanolamine solutions in the temperature range 75−140 °C. Tr. Gos. Nauchno-Issled. Proektn. Inst. Azotn. Promsti. 1959, 10, 54−82. (28) Aronu, U. E.; Gondal, S.; Hessen, E. T.; Haug-Warberg, T.; Hartono, A.; Hoff, K. A.; Svendsen, H. F. Solubility of CO2 in 15, 30, 45 and 60 mass % MEA from 40 to 120 °C and model representation using the extended UNIQUAC framework. Chem. Eng. Sci. 2011, 66, 6393−6406. (29) Brúder, P.; Lauritsen, K. G.; Mejdell, T.; Svendsen, H. F. CO2 capture into aqueous solutions of 3-methylaminopropylamine activated dimethyl-monoethanolamine. Chem. Eng. Sci. 2012, 75, 28−37. (30) Jou, F.-Y.; Mather, A. E.; Otto, F. D. The solubility of CO2 in a 30 mass percent monoethanolamine solution. Can. J. Chem. Eng. 1995, 73, 140−147. (31) Ma’mun, S.; Nilsen, R.; Svendsen, H. F.; Juliussen, O. Solubility of carbon dioxide in 30 mass % monoethanolamine and 50 mass % methyldiethanolamine solutions. J. Chem. Eng. Data 2005, 50, 630− 634. (32) Portugal, A.; Sousa, J.; Magalhaes, F.; Mendes, A. Solubility of carbon dioxide in aqueous solutions of amino acid salts. Chem. Eng. Sci. 2009, 64, 1993−2002. (33) Speyer, D.; Maurer, G. Solubility of hydrogen sulfide in aqueous solutions of piperazine in the low gas-loading region. J. Chem. Eng. Data 2011, 56, 763−767.

ASSOCIATED CONTENT

S Supporting Information *

Model details as described in the text; comparison between the measured and calculated CO2 species. This material is available free of charge via the Internet at http://pubs.acs.org.

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AUTHOR INFORMATION

Corresponding Author

*E-mail: [email protected]. Notes

The authors declare no competing financial interest.

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ACKNOWLEDGMENTS We thank Benedikt Ullmann for assistance in some measurements with the synthetic method. REFERENCES

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