Carbon dioxide solubility in aqueous potassium salt solutions of l -proline and dl -α-aminobutyric acid at high pressures

J. Chem. Thermodynamics 83 (2015) 110–116

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Carbon dioxide solubility in aqueous potassium salt solutions of L-proline and DL-a-aminobutyric acid at high pressures Yu-Tzung Chang a, Rhoda B. Leron a,b, Meng-Hui Li a,⇑ a b

R&D Center for Membrane Technology and Department of Chemical Engineering, Chung Yuan Christian University, Chung Li 32023, Taiwan, ROC School of Chemical Engineering and Chemistry, Mapúa Institute of Technology, Manila 1002, Philippines

a r t i c l e

i n f o

Article history: Received 9 September 2014 Received in revised form 8 December 2014 Accepted 10 December 2014 Available online 18 December 2014 Keywords: Amino acid salt Carbon dioxide solubility Kent–Eisenberg model

a b s t r a c t In the present work, the solubility of CO2 in aqueous solutions of potassium prolinate (KPr) and potassium a-aminobutyrate (KAABA) was measured at temperatures (313.2, 333.2, and 353.2) K and CO2 partial pressures up to 1000 kPa for amino acid salt concentrations: KPr, w = (7.5, 14.5, and 27.4 wt%) and KAABA, w = (6.9, 13.4, and 25.6 wt%). It was found that the CO2 absorption capacities of the studied amino acid salt systems were considerably high and comparable with that of industrially important alkanolamines including monoethanolamine. The CO2 loadings in aqueous potassium a-aminobutyrate at high pressures were also found to be generally higher than the loadings in aqueous potassium prolinate. A modified Kent–Eisenberg model was applied to correlate the CO2 solubility in the amino acid salt solution as function of CO2 partial pressure, temperature, and concentration. The model gave good representation of the (vapour + liquid) equilibrium data obtained for the amino acid salt systems studied, and provided accurate predictions of the solubility. Ó 2014 Elsevier Ltd. All rights reserved.

1. Introduction The leading technology today for post-combustion CO2 capture is based on the reversible chemical absorption of CO2 in aminebased solvents. The process most commonly utilises aqueous solutions of alkanolamines particularly monoethanolamine (MEA). The wide use of MEA as CO2 absorbent is due to its fast CO2 reaction kinetics, low solvent cost, high alkalinity, and other desirable characteristics. However, the use of MEA also entails a number of process limitations and other disadvantages. Aqueous MEA is corrosive at high concentrations, forms degradation products due to long exposure with CO2 and O2, and exhibits high vapour pressure causing significant solvent losses during processing. Also, it has a high heat of reaction with CO2, which leads to high energy requirement during solvent regeneration [1,2]. Thus, development of new and better solvents than aqueous MEA has always been necessary. Amino acid salts are proposed as suitable alternatives to aqueous alkanolamine solvents [3–6]. Due to the presence of amino functional groups in their molecules, amino acid salts, like alkanolamines, have fast reactivity and appreciable absorption capacity for CO2 [4,7–11]. The reaction kinetics of potassium salts of some ⇑ Corresponding author. Tel.: +886 3 265 4109; fax: +886 3 265 4199. E-mail address: [email protected] (M.-H. Li). http://dx.doi.org/10.1016/j.jct.2014.12.010 0021-9614/Ó 2014 Elsevier Ltd. All rights reserved.

amino acids were found to be even faster than MEA’s [12]. Additional advantages of aqueous amino acid salt solutions include low volatility and resistance to oxidative degradation [13]. They also have surface tension and viscosity that are similar to those of water, which makes their application more practical [14–17]. It has also been demonstrated that the utilisation of this type of solvents would lead to a more environment-friendly and energyeffective separation process for CO2 capture [18]. In the present work, we investigated the CO2 absorption capacities of two amino acid salts that are potential CO2 absorbents—the potassium salts of L-proline and DL-a-aminobutyric acid. Proline is a secondary amino acid with a distinctive cyclic structure, which includes an amine group while DL-a-aminobutyric acid is a sterically-hindered amino acid having an ethyl group substituent. The potassium salt of L-proline (KPr) is proposed by several authors as a candidate absorbent for CO2 capture due to its fast reaction rate and high absorption capacity for CO2 [10,12,19]. van Holst et al. [12] examined the absorption kinetics of CO2 in potassium salts of several amino acids including 6-aminohexanoic acid, b-alanine, Larginine, L-glutamic acid, DL-methionine, L-proline, and sarcosine at T = 298 K. They reported that the potassium salt of L-proline (and sarcosine) has relatively high apparent rate constant and low pKa making it a promising CO2 absorbent. In a more recent study, Paul and Thomsen [10], used a zwitterionic mechanism, which generally applies for the reaction of CO2 with primary and secondary

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Y.-T. Chang et al. / J. Chem. Thermodynamics 83 (2015) 110–116

alkanolamines, to describe the kinetics of reaction of CO2 with KPr. They found that the overall reaction rate constant for the latter is much higher than that of potassium threonate (KThr) and potassium taurate (KTau). The solvent has also been shown to have higher absorption capacity compared with the benchmark alkanolamine, MEA, and (vapour + liquid) equilibrium (VLE) data for low CO2 partial pressures (up to 30 kPa) have been reported [5,19]. On the other hand, the potassium salt of DL-a-aminobutyric acid (KAABA) has been found to have a significantly higher net cyclic capacity compared with MEA and potassium salts of some other amino acids [5]. The results of the study suggest that KAABA could be an energy-efficient alternative to MEA. However, except for the results of the (absorption + desorption) experiments at 15 kPa presented by Song et al. [5], no other VLE data has been reported for CO2 absorption in KAABA. The objective of this work was to provide VLE data for the absorption of CO2 in aqueous KPr and aqueous KAABA solutions at high pressures. The solubility of CO2 in the solvents was measured at temperatures (313.2, 333.2, and 353.2) K and pressures up to 1000 kPa. In addition, a modified Kent–Eisenberg model [20] was used to represent the correlation of the measured CO2 solubility (aCO2 , expressed as mole CO2 per mole amine) with CO2 partial pressure, temperature, and amino acid salt (AAS) concentration. 2. Experimental 2.1. Chemicals The aqueous potassium salt solutions of L-proline (Pr, mass fraction purity > 0.99), DL-a-aminobutyric acid (AABA, purity > 99 wt%), and sarcosine (Sr, mass fraction purity > 0.98) were prepared by adding an equimolar amount of KOH (pellets, mass fraction purity > 0.85) to the amino acid in a volumetric flask with deioniseddistilled water. The description and sources of all chemicals used, including those used in the validation experiments (Sr and MEA), are listed in table 1. They were used without further purification. The exact purity of KOH was determined by titration (1.0 M HCl, Scharlau) using a Metrohm 888 Titrando potentiometric titrator. The instrument allowed pH measurement to an accuracy of ±0.003 and was calibrated against standard buffer solutions before use. The exact molar concentrations of the amino acid salt solutions prepared were measured potentiometrically with standard HCl solution (1.0 M). The selected concentrations were (0.5, 1.0, and 2.0) M which had corresponding mass fraction compositions of (7.5, 14.5, and 27.4) wt% for KPr and (6.9, 13.4, and 25.6) wt% for KAABA. The standard uncertainty of the reported concentrations is ur(w) = ±0.005.

installed air bath. One is a 1.0-L vapour-recirculation equilibrium cell where the sample was contained while the other was a 0.5-L chamber connected to the equilibrium cell to increase the volume of the vapour phase. The temperature of the air bath was controlled by a basic immersion circulator provided by NESLAB Instruments (model EX-810B). The pressure was measured using a pressure transmitter (Druck PTX 1400), with an uncertainty of ±0.15% FS (range = 0 to 1013 kPa), connected to a pressure indicator (Vastech Scientific Co. Ltd.), which allowed reading to up to 0.1 kPa. Measurements were performed at pressures from (3 to 1000) kPa. In each experiment, the equilibrium cell was first purged with N2 or CO2 before the degassed sample (approximately 250 mL) was loaded. After loading the sample, the system was allowed to stabilize at the desired temperature. Then, CO2 was introduced until the system reaches the desired total pressure. For measurements involving CO2 partial pressures lower than 101.3 kPa, the CO2 stream was mixed with N2 as it was introduced into the equilibrium cell. The partial pressure of CO2 was determined by measuring the composition of the vapour phase. The composition (N2/CO2 ratio) of the vapour phase was monitored and determined through on-line gas chromatography (Shimadzu model GC-8A). It was assumed that the system was at equilibrium when the total pressure and gas phase composition remained constant for at least 2 h. Upon reaching equilibrium, a known weight of the liquid sample was withdrawn from the cell into a flask with excess NaOH (1.0 M) in order to convert free dissolved CO2 into non-volatile ionic species. Then, the CO2 loading in the sample was determined by potentiometric titration using the barium chloride method [23]. Sample analysis was repeated three to five times. The standard uncertainty (ur) of the measured loading is estimated to be ur(a) = ±0.03. 3. Results 3.1. Validation of experimental method To validate the experimental method used in this work, CO2 solubility in 15.3 wt% MEA solution at T = 313.2 K and 42.4 wt% potassium sarcosinate solution (KSr) at T = 353.2 K were measured. The VLE values obtained are given in tables 2 and 3 along with those reported by Lee et al. [24] and Park et al. [25] for CO2–MEA–H2O and Kang et al. [26] for CO2–KSr–H2O systems, respectively. Graphical representations of the results are also provided in figure 1 showing that the experimental values are in reasonable agreement with the literature data. Furthermore, for each system, the consistency of the results was examined by fitting an empirical polynomial function, equation (1), to the combined sets of data using the non-linear least squares method.

2.2. Measurement of CO2 solubility 1 2

1

The solubility of CO2 in each sample was measured using a (vapour + liquid) equilibrium setup, which we used and described in our previous works [21–23]. The setup consisted of two stainless steel cylinders, which were vertically mounted in a thermostat-

pCO2 =kPa ¼ A þ BðaCO2 =mol  mol Þ þ CðaCO2 =mol  mol Þ 1 3

1 4

þ DðaCO2 =mol  mol Þ þ EðaCO2 =mol  mol Þ 1 5

þ FðaCO2 =mol  mol Þ :

ð1Þ

TABLE 1 Provenance and mass fraction purity of chemicals used in the study. Chemical

CAS No.

Mass fraction purity)

Source

Method of purification

Monoethanolamine (MEA) Sarcosine (Sr)

141-43-5 107-97-1 147-85-3

>0.99 >0.98 >0.99

Tedia Co. Alfa Aesar Alfa Aesar

None None None

2835-81-6

>0.99

Alfa Aesar

None

1310-58-3 124-38-9

>0.99 >0.9999

Sigma–Aldrich Ming Yang Special Gas Co., Ltd.

None None

L-Proline DL-

(Pr)

a-aminobutyric acid (AABA)

Potassium hydroxide (KOH) Carbon dioxide (CO2)

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Y.-T. Chang et al. / J. Chem. Thermodynamics 83 (2015) 110–116

TABLE 2 Solubility of CO2 in aqueous (w = 15.3 wt%) MEA solution at T = 313.2 K.a Park et al. [25] b

pCO2 =kPa

aCO2

3.9 19.9 132.2 137.9 245.4 365.0 446.0 929.0

0.502 0.562 0.684 0.678 0.748 0.800 0.797 0.808

This work

pCO2 =kPa

aCO2

pCO2 =kPa

aCO2

5.88 16.48 71.49 484.7

0.505 0.562 0.636 0.801

7.0c 7.2 14.6 19.1 32.0 59.9 93.8 152.8 240.7 362.3 462.7 562.2 762.4 846.9 994.4 AAD/% = 12.5

0.503 0.511 0.557 0.561 0.608 0.651 0.678 0.738 0.759 0.799 0.827 0.837 0.859 0.883 0.894

1000

pCO2 / kPa

Lee et al. [24]

10000

100

10

a

Standard uncertainties: u(T) = ±0.1 K; ur(pCO2 ) = ±1.5  103; ur(aCO2 ) = ±0.03; ur(w) = ±0.005. b aCO2 is the CO2 loading of the solvent in mol CO2 per mol amine. c Data in italics were obtained using mixed (CO2 + N2) gas stream.

TABLE 3 Solubility of CO2 in aqueous (w = 42.4 wt%) potassium sarcosinate (KSr) solution at T = 353.2 K.a Kang et al. [26]

1 0.2

0.4

0.6

0.8

1.0

−1 αCO2 / (molCO2⋅ molsolvent )

FIGURE 1. CO2 solubility in aqueous (w = 15.3 wt%) MEA solution at T = 313.2 K: (j, this work; h, Lee et al. [24]; 4, Park et al. [25]); and in aqueous (w = 42.4 wt%) KSr solution at T = 353.2 K: (d, this work, s, Kang et al. [26]); lines, calculated using equation (1).

This work

pCO2 =kPa

aCO2 b

pCO2 =kPa

aCO2

8.8 24.6 68.6 153.9 281.1 429.4 577.8 731.5 938.4

0.475 0.543 0.607 0.667 0.717 0.756 0.785 0.811 0.843

4.8c 11.6 50.7 77.9 147.9 273.1 322.9 479.1 614.5 762.1 875.9 962.3 AAD/% = 14.5

0.374 0.489 0.564 0.627 0.657 0.726 0.739 0.748 0.762 0.803 0.824 0.837

a

Standard uncertainties: u(T) = ±0.1; ur(pCO2 ) = ±0.0015; ur(aCO2 ) = ±0.03; K; ur(w) = ±0.005. b aCO2 is the CO2 loading of the solvent in mol CO2 per mol amino acid salt. c Data in italics were obtained using mixed (CO2 + N2) gas stream.

Here pCO2 is the partial pressure of CO2 in kPa and aCO2 is the CO2 loading in the sample in mol CO2 per mol solvent. A to F are empirical constants obtained from the regression whose values are listed in table 4. The average absolute deviations (AADs) between our experimental and the predicted values by equation (1) were (12.5 and 14.5)% for CO2–MEA–H2O and CO2–KSr–H2O systems, respectively. The AAD is defined as

, N   1X   AAD ¼ Pexp  Pcal Pexp ; N i

ð2Þ

where Pexp and Pcal are the experimental and the predicted values, respectively, and N is the number of data points. 3.2. Solubility of CO2 in aqueous amino acid salt solutions The measured CO2 solubility for aqueous potassium prolinate (KPr) and aqueous potassium a-aminobutyrate (KAABA) systems are given in tables 5 and 6, respectively. Figures 2 and 3 also present the VLE results for the two systems for different molar

TABLE 4 Empirical constants of equation (1). Constant

CO2–MEA–H2O

CO2–KSr–H2O

A B C D E F AAD/%

8898 67,297 200,470 292,020 204,570 52,575 12.5

36,448 327,896 1,154,535 1,988,666 1,673,569 548,144 14.5

concentrations at the temperatures studied. It can be observed that the CO2 loading increases with increasing partial pressure of CO2 but decreases with increasing molar concentration of the amino acid salt solution. Also, for a given molar concentration and CO2 partial pressure, the CO2 loading decreases systematically as temperature increases. This behaviour can be attributed to the exothermic nature of CO2 absorption. It should also be noted that in the low CO2 partial pressure range, for a given molar concentration and absorption temperature, CO2 loading in the KPr solution is higher than in the KAABA solution. However, the loading in the KAABA solution reaches a higher value (than for KPr) at high CO2 partial pressure. As shown in figure 4, such effect of the CO2 partial pressure on CO2 loading becomes more obvious as the molar concentration of the absorbent increases. The high CO2 loadings in KAABA solutions can be attributed to the presence of unstable carbamates in the solvent during CO2 absorption. Such instability of the carbamates formed is caused by the steric hindrance created by the a-substituent (–CH2CH3) present in the amino acid [27]. Hence, unlike with KPr where stable carbamates are formed with CO2, the reaction of CO2 with KAABA favours bicarbonate formation resulting in the release of free amines and greater CO2 loading [13,28]. Figure 5 depicts the comparison of the absorption capacities of the studied amino acid salt systems at T = 313.2 K with known alkanolamine absorbents and other amino acid salt systems. As shown, at the same temperature and CO2 partial pressure, the

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Y.-T. Chang et al. / J. Chem. Thermodynamics 83 (2015) 110–116 TABLE 5 Solubility of CO2 in aqueous potassium prolinate (KPr)a solutions. w = 7.5 wt% KPr

w = 14.5 wt% KPr

w = 27.4 wt% KPr

w = 6.9 wt% KAABA

pCO2 =kPa

10.2b 67.9 130.3 246.5 374.6 429.0 532.7 711.8 928.6

0.821 0.984 1.026 1.055 1.092 1.101 1.110 1.136 1.160

T = 313.2 K 2.9 0.532 20.7 0.643 68.4 0.737 130.6 0.769 230.5 0.829 463.3 0.868 600.5 0.893 721.1 0.909 837.2 0.920

8.3 29.7 64.8 179.2 237.8 308.3 549.6 708.1 824.9

0.351 0.404 0.436 0.490 0.520 0.542 0.563 0.572 0.579

9.9b 55.2 121.7 241.1 378.6 566.2 616.3 828.3 942.7

0.722 0.883 0.921 0.976 1.011 1.067 1.080 1.105 1.124

10.7 71.9 134.5 229.2 448.1 606.3 814.8 977.6

T = 313.2 K 0.567 0.759 0.814 0.876 0.921 0.945 0.978 1.004

4.2 19.3 57.2 154.3 277.6 462.1 549.5 719.2 844.3 927.3

0.622 0.726 0.861 0.940 0.971 1.016 1.024 1.046 1.057 1.064

T = 333.2 K 9.6 0.514 57.6 0.622 88.5 0.651 155.9 0.699 305.7 0.738 426.8 0.763 517.6 0.789 714.5 0.809 821.2 0.821

5.9 16.8 26.0 159.7 338.3 460.3 577.1 706.3 811.3

0.287 0.340 0.357 0.420 0.454 0.469 0.481 0.492 0.501

7.8 49.7 77.9 107.2 327.6 592.1 716.8 922.9

0.632 0.812 0.829 0.842 0.928 0.991 1.010 1.040

4.6 32.1 97.6 274.2 347.2 511.7 638.9 842.3 921.4

T = 333.2 K 0.406 0.614 0.705 0.772 0.820 0.870 0.890 0.900 0.910

6.5 87.6 188.1 247.2 362.5 438.7 627.4 780.1 955.3

0.521 0.752 0.830 0.854 0.874 0.915 0.941 0.959 0.986

8.1 18.3 47.8 77.6 135.2 270.0 403.3 657.9 847.3

T = 353.2 K 0.322 0.433 0.548 0.588 0.633 0.709 0.759 0.798 0.831

0.571 0.662 0.814 0.874 0.916 0.930 0.948 0.965 0.979 0.985

3.0 24.1 54.5 159.7 211.1 384.3 476.9 600.1 826.0

pCO2 =kPa

0.242 0.291 0.323 0.365 0.387 0.401 0.424 0.452 0.466

a

Standard uncertainties: u(T) = ±0.1 K; ur(pCO2 ) = ±1.5  103; ur(aCO2 ) = ±0.03; ur(w) = ±0.005. b Data in italics were obtained using mixed (CO2 + N2) gas stream.

CO2 loadings in 14.5 wt% KPr and 13.4 wt% KAABA solutions are higher than those in 15.3 wt% MEA. Although the absorption capacities of the systems studied were found inferior compared to that of 1.0 M potassium glycinate (KGly), their performances were comparable with those of KThr and KTau. It should also be noted that the absorption capacity of the KAABA solution was comparable with those of 30 wt% MDEA and 30 wt% AMP solutions. Such result can also be attributed to the bulky functional group of AABA [5]. It can also be noted that the VLE curves for KPr and KAABA follow typical behaviour of primary or secondary (as in MEA) and sterically hindered or tertiary (as in AMP or MDEA) alkanolamines, respectively. From the comparison presented, it can be said that on the basis of absorption capacity, amino acid salt solutions may be promising alternatives to alkanolamine absorbents. 4. Modelling In this work, we used a modified form of the Kent–Eisenberg model proposed by Li and Shen [20] to correlate the CO2 loading in the aqueous amino acid salt solution with temperature, CO2 partial pressure, and AAS concentration. This model allows, with simplicity, the calculation of CO2 solubility in amine solvents using equilibrium constant expressions derived from equilibrium reactions, which are assumed to occur between CO2 and the amine. It has been shown that the absorption of CO2 in aqueous amino acid salt solutions follow a mechanism similar to that of aqueous amine solutions [28]. Hence, the equilibrium reactions can be described accordingly as follows (for potassium prolinate): 

K1

 PrCOO þ H2 O !  Pr þ HCO3 ;

ð3Þ

aCO2

pCO2 =kPa

w = 25.6 wt% KAABA

aCO2

pCO2 =kPa

aCO2

22.7 57.2 124.3 335.4 533.2 732.7 822.1 914.3

0.479 0.558 0.670 0.715 0.768 0.805 0.821 0.844

88.1 173.2 383.9 498.6 685.3 767.2 977.3

0.501 0.580 0.675 0.697 0.739 0.746 0.762

74.2 119.3 293.6 438.6 550.9 640.2 790.1 918.3

0.387 0.456 0.550 0.582 0.617 0.627 0.660 0.680

a

Standard uncertainties: u(T) = ±0.1 K; ur(pCO2 ) = ±1.5  103; ur(aCO2 ) = ±0.03; ur(w) = ±0.005. b Data in italics were obtained using mixed (CO2 + N2) gas stream.

10000

1000

pCO2 / kPa

7.1 16.8 87.6 199.4 346.7 422.8 564.3 644.9 751.1 845.9

T = 353.2 K 7.0 0.411 17.0 0.475 65.8 0.553 161.0 0.613 340.2 0.681 506.2 0.713 673.2 0.738 764.9 0.754 867.2 0.766

pCO2 =kPa

aCO2

w = 13.4 wt% KAABA

aCO2

pCO2 =kPa

aCO2

TABLE 6 Solubility of CO2 in aqueous potassium aminobutyrate (KAABA)a solutions.

100

10

1 0.0

0.2

0.4

0.6

0.8

1.0

1.2

−1

α / (molCO2 ⋅ molAAS ) FIGURE 2. CO2 solubility in aqueous KPr solutions at w = 7.5 wt%: (h, T = 313.2 K; s, T = 333.2 K; 4, T = 353.2 K); w = 14.5 wt%: (j, T = 313.2 K; d, T = 333.2 K; N, T = 353.2 K); and w = 27.4 wt% (5, T = 313.2 K; }, T = 333.2 K; , T = 353.2 K); lines, calculated using the modified Kent–Eisenberg model.

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Y.-T. Chang et al. / J. Chem. Thermodynamics 83 (2015) 110–116

10000

10000

1000

1000

pCO2 / kPa

pCO2 / kPa

100

100

10

10

1

0.1

1 0.0

0.2

0.4

0.6

0.8

1.0

0.0

1.2

0.3

0.6

0.9 −1 amine

α / (molCO2⋅ mol

−1 αCO2 / (molCO2 ⋅ molAAS )

FIGURE 3. CO2 solubility in aqueous KAABA solutions at w = 6.9 wt%: (h, T = 313.2 K; s, T = 333.2 K; 4, T = 353.2 K); w = 13.4 wt%: (j, T = 313.2 K; d, T = 333.2 K; N, T = 353.2 K); and w = 25.6 wt% (5, T = 313.2 K; }, T = 333.2 K; , T = 353.2 K); lines, calculated using the modified Kent–Eisenberg model.

1.2

)

FIGURE 5. Comparison of CO2 solubility in aqueous amine absorbents at T = 313.2 K: j, 14.5 wt% KPr; d, 13.4 wt% KAABA (experimental data from this work); h, 15.3 wt% MEA [23]; s, 30 wt% MDEA [23]; , 30 wt% AMP [30]; 5, 1.0 M KGly [4]; }, 1.0 M KThr [4]; 4, 1.0 M KTau [31]; —, calculated using the modified Kent–Eisenberg model; ———, smoothed lines.

10000

where Pr represents the prolinate ion. The expressions for the apparent equilibrium constants, Ki, for the above equations can be expressed as:

pCO2 / kPa

1000

K1 ¼

½Pr ½HCO3  ; ½ PrCOO 

ð8Þ

K2 ¼

½Pr ½Hþ  ; ½ PrHþ 

ð9Þ

K3 ¼

½Hþ ½HCO3  ; ½CO2 

100

K 4 ¼ ½Hþ ½OH ;

10

K5 ¼ 1 0.0

0.2

0.4

0.6

0.8

1.0

pCO2 ¼ HCO2 ½CO2 :

FIGURE 4. Comparison of CO2 solubility in the studied amino acid salt systems at T = 333.2 K: (KPr: h, 7.5 wt%; s, 14.5 wt%; 4, 27.4 wt%; (KAABA: j, 6.9 wt%; d, 13.4 wt%; N, 25.6 wt%); lines, calculated using the modified Kent–Eisenberg model.

þ K2



þ

PrH $ Pr þ H ; K3

þ

H2 O þ CO2 $ H

þHCO3 ;

K4

H2 O $ Hþ þOH ; K5

HCO3 $ Hþ þ CO2 3 ;

ð11Þ ð12Þ

The partial pressure of CO2 is related to the dissolved (physically) CO2 in the solvent by Henry’s law as in equation (13).

1.2

−1 ) αCO2 / (molCO2 ⋅ molAAS



½Hþ ½CO2 3  : ½HCO3 

ð10Þ

ð13Þ

Here, pCO2 is the CO2 partial pressure in kPa and HCO2 is the Henry’s law constant. The mass balances between the reacting species are described by the following equations:

mAAS ¼ ½Pr  þ ½ PrCOO  þ ½ PrHþ ;

ð14Þ

mAAS aCO2 ¼ ½CO2  þ ½ PrCOO  þ ½HCO3  þ ½CO2 3 ;

ð15Þ

ð4Þ ð5Þ ð6Þ ð7Þ

where mAAS is the molar concentrations of the amino acid salt in kmol  m3 and aCO2 is the CO2 loading in the solvent in mol CO2 per mol amino acid salt. The charge balance among the species is:  ½Kþ  þ ½Hþ  ¼ ½Pr  þ 2½ PrCOO  þ ½HCO3  þ 2½CO2 3  þ ½OH :

ð16Þ

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Y.-T. Chang et al. / J. Chem. Thermodynamics 83 (2015) 110–116

For potassium a-aminobutyrate, similar expressions can be used to describe the equilibrium reactions with exception of the carbamate hydrolysis reaction (equation (3)) since sterically-hindered amines (or amino acid salts, e.g. a-aminobutyrate) do not form stable carbamates upon reaction with CO2. To describe the amine protonation reaction for potassium a-aminobutyrate, equation (4) can be re-written as: K2

RNHþ3 $ RNH2 þ Hþ ;

ð17Þ

where R„ –CH(CH2CH3)COO. For the apparent equilibrium constants K3 to K5 and the Henry’s law constant, results obtained by Kent and Eisenberg [29] were used. The expressions are as follows:

h K 3 =ðkmol  m3 Þ ¼ exp 241:818 þ 29:8253  104 =ðT=KÞ 8

2

1000

Predicted pCO2 / kPa



10000

100

10 3

11

 1:48528  10 =ðT=KÞ þ 0:332647  10 =ðT=KÞ i  0:282393  1013 =ðT=KÞ4 ; ð18Þ 1

h 2 K 4 =ðkmol  m3 Þ ¼ exp 39:5554  9:879  104 =ðT=KÞ þ 0:568827  108 =ðT=KÞ2 i  0:146451  1011 =ðT=KÞ3 þ 0:136145  1013 =ðT=KÞ4 ; ð19Þ

h K 5 =ðkmol  m3 Þ ¼ exp 294:74 þ 36:4385  104 =ðT=KÞ

1

10

100

1000

10000

Measured pCO2 / kPa FIGURE 6. Parity plot of predicted vs. measured CO2 partial pressures over various aqueous KPr solutions: h, 7.5 wt%; s, 14.5 wt%; 4, 27.4 wt%.

 1:84157  108 =ðT=KÞ2 þ 0:415792  1011 =ðT=KÞ3 i  0:354291  1013 =ðT=KÞ4 ; ð20Þ

10000

h i h 1 ¼ exp 22:2819  1:3831  104 =ðT=KÞ HCO2 = mmHg  ðkmol  m3 Þ þ 0:06914  108 =ðT=KÞ2  0:0156  1011 =ðT=KÞ3 i ð21Þ þ 0:012  1013 =ðT=KÞ4 :

n K 1KPr =ðkmol  m3 Þ ¼ exp 22:245  3869:0=ðT=KÞ

Predicted pCO2 / kPa

Using the above equations, the apparent equilibrium constant (K1 and K2) expressions for the main amine reactions described in equations (3) and (4) for KPr and equation (17) for KAABA were determined. Experimental CO2 partial pressures and loadings were used and empirical coefficients of the said expressions were regressed by the least squares method. The resulting equations are:

1000

10

 370270=ðT=KÞ2  29:222aCO2  o  14:393a2CO2 þ 4:7457mAAS = kmol  m3 ;

ð22Þ

n

K 2 =ðkmol  m3 Þ ¼ exp 11:601 þ 6049:2=ðT=KÞ  722890=ðT=KÞ2

1

2 CO2

 25:001aCO2 þ 3:5124a  o  3:7132mAAS = kmol  m3 ; 

K 1KAABA = kmol  m

3



10

100

1000

10000

Measured pCO2 / kPa FIGURE 7. Parity plot of predicted vs. measured CO2 partial pressures over various aqueous KAABA solutions: h, 6.9 wt%; s, 13.4 wt%; 4, 25.6 wt%.

¼ exp 12:025  996:47=ðT=KÞ  490220=ðT=KÞ2  9:3856aCO2 þ 11:304a

1

ð23Þ

n

2 CO2

100

o þ 1:3918mAAS =ðkmol  m Þ : 3

ð24Þ

The predicted CO2 partial pressures from the model are presented as a function of CO2 loading, together with the experimental results in figures 2 and 3 for aqueous potassium prolinate and potassium a-aminobutyrate systems, respectively. In figures 6 and 7, parity plots of the predicted and measured CO2 partial pressures for all studied conditions are presented. It can be deduced that the predicted and the measured values agree reasonably well. The calculated values of AAD for KPr and KAABA systems are (13.1 and 15.0)%, respectively. We also performed a comparison of the predicted values from the model for potassium prolinate with the experimental results reported by Majchrowicz and Brilman

[19] at low CO2 partial pressures (3.7 to 30.0) kPa. The predicted values were found to be consistently lower than the literature data with a maximum deviation of 12.1%. It can be argued that such deviation is well within the reported accuracy of the model proposed in the present work. These results suggest that the applied model, along with the regressed coefficients, satisfactorily correlated the CO2 partial pressure above the solvent systems with temperature, CO2 loading and amino salt concentration for all experimental conditions considered in this work. 5. Conclusions The CO2 solubility in aqueous potassium salts of L-proline and

a-aminobutyric acid were measured at temperatures (313.2,

DL-

116

Y.-T. Chang et al. / J. Chem. Thermodynamics 83 (2015) 110–116

333.2, and 353.2) K and CO2 partial pressures up to 1000 kPa for amino acid salt concentrations (0.5, 1.0, and 2.0) M. Results show that CO2 loadings in the solvents increase with increasing CO2 partial pressure and decrease as temperature increases. It was also found that the absorption capacities of the amino acid salt systems studied are in the same order of magnitude as known alkanolamines’ such as MEA, AMP, and MDEA. A modified Kent–Eisenberg model was successfully applied to represent the CO2 solubility (in terms of CO2 loading) as a function of CO2 partial pressure, temperature, and amino acid salt concentration. The respective average absolute deviations of the predicted results (as CO2 partial pressure) from the experimental values are (13.1 and 15.0)% for potassium prolinate and potassium a-aminobutyrate systems. Hence, it can be said that the model applied would be useful in the prediction of (vapour + liquid) equilibrium for the systems studied that are important in the design and optimisation of a suitable CO2 absorption process. Acknowledgement This research was supported by Grant MOST-103-2221-E-033067-MY3 from the Ministry of Science and Technology of the Republic of China. We also thank Chung Yuan Christian University for financial support. References [1] A.L. Kohl, R.B. Nielsen, Gas Purification, fifth ed., Gulf Publishing Company, Houston, TX, 1997. [2] R. Idem, P. Tontiwachwuthikul, Ind. Eng. Chem. Res. 45 (2006). 2413-2413. [3] P.H.M. Feron, N. ten Asbroek, E.S. Rubin, D.W. Keith, C.F. Gilboy, M. Wilson, T. Morris, J. Gale, K. Thambimuthu, New Solvents based on Amino-Acid Salts for CO2 Capture from Flue Gases, Greenhouse Gas Control Technologies, vol. 7, Elsevier Science Ltd, Oxford, 2005. pp. 1153–1158. [4] A.F. Portugal, J.M. Sousa, F.D. Magalhaes, A. Mendes, Chem. Eng. Sci. 64 (2009) 1993–2002. [5] H.-J. Song, S. Park, H. Kim, A. Gaur, J.-W. Park, S.-J. Lee, Int. J. Greenhouse Gas Control 11 (2012) 64–72.

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JCT 14-506