Article pubs.acs.org/jced
CO2 Solubility Measurements and Modeling for Tertiary Diamines Hiroshi Machida,*,† Hidetaka Yamada,‡ Yuichi Fujioka,§ and Shin Yamamoto‡ †
Department of Chemical Engineering, Graduate School of Engineering, Nagoya University, Furo-cho, Chikusa-ku, Nagoya 464-8601, Japan ‡ Chemical Research Group, Research Institute of Innovative Technology for the Earth (RITE), 9-2 Kizugawadai, Kizugawa-shi, Kyoto, 619-0292, Japan § Department of Environmental Sciences, International College of Arts and Sciences, Fukuoka Women’s University, 1-1-1 Kasumigaoka, Higashi-ku, Fukuoka, 813-8529, Japan ABSTRACT: CO2 solubility in the solutions of three types of tertiary diamines, N,N,N′,N′-tetramethylethylenediamine, N,N,N′,N′-tetramethyl-1,3-diaminopropane, and N,N,N′,N′-tetramethyl-1,6-diaminohexane, was measured using vapor−liquid equilibrium equipment at temperatures from (313 to 393) K and pressures from (10 to 500) kPa. For these tertiary diamines, the CO2 solubility was calculated based on a simplified chemical equilibrium model. It was observed that alkyl chain length between two amines has an important role in CO2 solubility. To clarify this role, we also performed quantum mechanical calculations using the density functional theory and a continuum solvation model.
1. INTRODUCTION
The tertiary amine is suited for high-pressure process gas, as it has a low heat of absorption and high CO2 solubility at high pressures of (0.5 to 1) MPa. N-Methyldiethanolamine (MDEA) is widely used for the removal of carbon dioxide with modifiers (usually primary or secondary amines).9−20 Numerous studies on chemical absorption phenomena with MDEA have been reported.21−24 Apart from MDEA, we studied CO2 absorption using new types of tertiary amine solutions such as morphorin and imidazole at high pressure.25−27 The tertiary amine is considered to be useful also at low pressure separation process especially when it is blended with primary or secondary amines. We investigated the effects of substituents around the amino group on CO2 absorption-regeneration performance using 24 tertiary amine absorbents below atmospheric pressure.28 In this study, we extended the examination of tertiary amine absorbents to diamines, which contain two amino groups in one molecule and could be highly efficient to improve CO2 solubility as additives. We focused on the effect of alkyl chain length between the two amino groups and measured CO2 solubility in the aqueous solutions of three types of tertiary diamines: N,N,N′,N′-tetramethylethylenediamine (C2-diamine), N,N,N′,N′-tetramethyl-1,3-diaminopropane (C3-diamine), N,N,N′,N′-tetramethyl-1,6-diaminohexane (C6-diamine) using vapor−liquid equilibrium (VLE) equipment at temperatures from (313 to 393) K and pressures from (10 to
CO2 capture and storage (CCS) technology is expected to be an effective technology for the reduction of greenhouse gas emissions. Because CO2 capture requires a considerable amount of energy, it covers large portion of the total CCS cost. Therefore, capture technologies such as amine scrubbing, pressure swing adsorption, and membrane separation have been extensively studied to reduce the energy requirements.1−3 Amine scrubbing has been developed on an industrial scale and is the most feasible for CO2 large point sources.4 In this process, an aqueous amine solution contacts the process gas and dissolves CO2 selectively at an absorption column. Then, the CO2-loaded amine solution flows to a stripping column for CO2 stripping by heating, and the regenerated amine solution is reused at the absorption column. The energy requirement for CO2 capture and regeneration significantly depends on the amine chemical structure and process design. Key points for the development of high-efficiency amine absorbents with reduced energy cost include the following: (1) low heat of absorption, (2) large difference in CO2 solubility between absorption (313 K) and stripping (∼393 K) temperatures, (3) high CO2 absorption rate, and (4) high CO2 stripping rate. Several amine components and their mixtures have been studied to develop energy-saving chemical absorbents.5,6 Singh et al. reported about the structure dependence for CO2 solubility of various primary amines including diamines.7,8 The results for diamines showed that the absorption capacity was increased with increasing chain length between the amino groups in most cases. © XXXX American Chemical Society
Received: October 7, 2014 Accepted: January 27, 2015
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an aqueous amine sample solution (30 wt % amine), and the solution was stirred. The temperature was controlled with a heater. CO2 and N2 flowed into the equilibrium cell after passing through the humidifier at the same temperature as the cell. The condenser after the cell was set at 278 K using a chiller unit. The CO2 concentration in the gas phase was monitored by the CO2 analyzer. When equilibrium was reached, a small amount of solution was collected in a bottle (1 cc). CO2 solubility in the solution was determined by a total organic carbon analyzer (TOC-VCSH, Shimadzu). Furthermore, pKa values were measured with a titration apparatus (DL58 Titrator, Mettler Toledo) to examine the relationship between CO2 solubility and amine basicity. 2.3. CO2 Solubility Model. The CO2 solubility model was developed using simple chemical reaction equilibrium. In this model, activity coefficients of each component were assumed to be 1, and hydroxide (OH−) and carbonate (CO2− 3 ) were not considered since their concentrations will be negligibly small. Similar assumptions were also used in other CO2 solubility models.23,24 Although Tomizaki et al. showed that there is physical and chemical absorption of CO2,25 the former is negligible under 1 MPa. The following equation was developed using two chemical reaction constants. For chemical absorption for tertiary diamines, CO2 is loaded according to the following equation:
500) kPa. The amine concentration was set to 30 wt % since many researches for amine scrubbing used this concentration, and it is easy to compare or discuss about the absorption phenomena. pKa was measured with a titration apparatus and calculated using density functional theory (DFT) to discuss the relationship between amine basicity and CO2 solubility. A CO2 solubility model was developed considering the chemical equilibrium constant.
2. EXPERIMENTAL AND CALCULATION METHODS 2.1. Materials. Table 1 lists the purity and suppliers of the three types of tertiary diamines and CO2 and N2 used in this study. Table 1. Properties and Suppliers of Materials
diamine + CO2 + H 2O ↔ (diamineH+)(HCO−3 )
(1)
(diamineH+)(HCO−3 ) + CO2 + H 2O ↔ (HCO−3 )(H+diamineH+)(HCO−3 )
(2)
The equilibrium constants K1′ and K2′ are defined as follows: K1′ =
[(diamineH+)(HCO−3 )] [diamine][CO2 ]
(3)
K 2′ =
[(HCO−3 )(H+diamineH+)(HCO−3 )] [(diamineH+)(HCO−3 )][CO2 ]
(4)
Considering that the protonated amine and the bicarbonate HCO3− are one molecule and CO2 molecule concentration in liquid phase is proportional to the CO2 pressure, eqs 3 and 4 can be rewritten as Langmuir form:
2.2. CO2 Solubility Measurements. The VLE equipment is shown in Figure 1. The experimental procedure has been described in previous studies.29,30 CO2 solubility data obtained for aqueous amine solutions using this equipment agreed well with the literature data.30 The equipment comprised a CO2 flow meter, a N2 flow meter, an equilibrium cell, a condenser, a back pressure regulator, and an infrared CO2 analyzer (VA3001, Horiba). The equilibrium cell (500 mL) was filled with
0 [diamine] = Xdiamine (1 − α1)
(5)
0 [(diamine)(HCO−3 )] = Xdiamine α1
(6)
pCO = H[CO2 ]
(7)
2
K1′ =
α1 (1 − α1) ·pCO /H
(8)
2
K1 = K1′/H =
α1 (1 − α1) ·pCO
(9)
2
0 [(diamineH+)(HCO−3 )] = Xdiamine α1· (1 − α2)
(10)
0 [(HCO−3 )(H+diamineH+)(HCO−3 )] = Xdiamine α1·α2
(11)
K 2′ = Figure 1. Equilibrium equipment.
α2 (1 − α2) ·pCO /H 2
B
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Figure 2. CO2 solubility data for 30 wt % amine of (a) C2-diamine, (b) C3-diamine, (c) C6-diamine at ⧫, 313 K; ■, 343 K; correlation with eqs 14 and 15.
K 2 = K 2′/H =
α2 (1 − α2) ·pCO
where X0diamine is the initial amine concentration, pCO2 is partial pressure of CO2, H/MPa is Henry’s constant, and K1/(MPa−1) and K2/(MPa−1) are the combined parameters of K′ and H. α1 and α2 are CO2 solubility (molCO2/molamine) corresponding to the forms of (diamine)(HCO 3 − ) and (HCO 3 − )(H+diamineH+)(HCO3−), respectively. The CO2 solubility for diamine α is finally determined with the following equation: α = α1 + α2 =
pCO ·K1 2
1 + pCO ·K1 2
+
pCO ·K 2
ln K i = ln
1 + pCO ·K 2 2
ΔHi ⎛ 1 1⎞ − ⎜ − ⎟ R ⎝ T2 T1 ⎠ −1
1 n
∑ |αcalc − αexp|
diamineH+ + H3O+ ↔ H+diamineH+ + H 2O
(18)
pK a1 = log10
[diamine][H3O+] [diamineH+]
(19)
pK a2 = log10
[diamineH+][H3O+] [H+diamineH+]
(20)
where i = 1, 2 −1
where R/(J·K ·mol ) is a gas constant and ΔH/(kJ·mol ) is the reaction enthalpy. Parameters K and ΔH were determined with experimental data using a generalized reduced graduent (GRG) method. The objective function to determine the equilibrium constant and the reaction enthalpy is as follows. OF =
(17)
For these two reactions, the corresponding pKa values are represented by eqs 19 and 20.
(14)
(15) −1
diamine + H3O+ ↔ diamineH+ + H 2O
2
The temperature dependence for constant K is given by the van’t Hoff equation: K i313
393 K and
2.4. DFT Calculations. DFT calculations were performed at the B3LYP/6-311++G(d,p) level of theory in vacuum and aqueous phases to optimize the geometries of diamine and its protonated cations. For the aqueous phase optimizations, we employed the latest continuum solvation model (SMD) using the integral-equation-formalism polarizable continuum model protocol for bulk electrostatics.31,32 In an aqueous solution, tertiary diamine reacts with H+ according to the following equations:
(13)
2
▲,
From the relation between the reaction free energy and the equilibrium constant, the following equation is given
pK a = γ ΔE + β
(21)
where ΔE is the energy difference between diamine and diamine H+ (pKa1) or between diamine H+ and H+diamine H+ (pKa2) calculated from the DFT optimizations. Parameters γ
(16) C
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Table 2. Experimental Values of CO2 Solubility Data for the Diamines at Temperature T and CO2 Partial Pressure pCO2a,b
a
T/K
ptotal/kPa
CO2 vol % in vapor phase
pwater/kPa
313.2 313.2 313.2 313.2 313.2 313.2 343.2 343.2 343.2 343.2 343.2 343.2 393.2 393.2 393.2 393.2 393.2 393.2
132.2 142.3 151.9 227.6 227.0 431.7 141.5 162.2 171.9 241.2 241.5 451.9 301.6 322.4 342.0 401.6 402.3 611.8
9.9 19.8 39.6 49.8 100.0 100.0 9.9 19.8 39.6 49.8 100.0 100.0 5.0 9.9 19.8 39.6 49.8 100.0
6.9 6.9 6.9 6.9 6.9 6.9 29.3 29.3 29.3 29.3 29.3 29.3 186.3 186.3 186.3 186.3 186.3 186.3
313.2 313.2 313.2 313.2 313.2 313.2 343.2 343.2 343.2 343.2 343.2 343.2 393.2 393.2 393.2 393.2 393.2 393.2
130.9 140.4 158.6 224.8 225.2 435.9 141.5 161.9 170.8 245.3 248.5 454.9 301.0 324.0 341.4 404.8 401.8 612.8
9.9 19.8 39.6 49.8 100.0 100.0 9.9 19.8 39.8 49.9 100.0 100.0 9.8 19.8 39.6 49.8 100.0 100.0
7.0 7.0 7.0 7.0 7.0 7.0 29.5 29.5 29.5 29.5 29.5 29.5 187.6 187.6 187.6 187.6 187.6 187.6
313.2 313.2 313.2 313.2 313.2 313.2 343.2 343.2 343.2 343.2 343.2 343.2 393.2 393.2 393.2 393.2 393.2 393.2
130.6 140.4 158.2 226.9 228.6 438.5 140.1 160.0 169.7 255.0 254.8 456.1 301.3 320.7 342.0 405.1 402.0 614.9
10.1 19.8 39.7 49.8 100.0 100.0 9.9 19.7 39.6 49.8 100.0 100.0 9.9 19.8 39.6 49.8 100.0 100.0
7.1 7.1 7.1 7.1 7.1 7.1 29.9 29.9 29.9 29.9 29.9 29.9 190.2 190.2 190.2 190.2 190.2 190.2
pCO2/kPa
CO2 solubility/(molCO2/molamine)
C2-Diamine 30 wt % 12.4 26.8 57.5 109.8 220.1 424.8 11.1 26.3 56.5 105.5 212.2 422.6 11.4 26.9 61.7 107.2 216.0 425.5 C3-Diamine 30 wt % 12.3 26.4 60.1 108.4 218.2 429.0 11.1 26.2 56.3 107.7 219.0 425.4 11.2 26.9 61.0 108.1 214.2 425.2 C6-Diamine 30 wt % 12.4 26.4 60.0 109.5 221.6 431.5 10.9 25.6 55.4 112.0 225.0 426.2 11.0 25.8 60.2 107.0 211.8 424.7
CO2 solubility (mass fraction)
0.644 0.781 0.881 0.949 1.01 1.068 0.208 0.357 0.549 0.701 0.817 0.905 0.023 0.041 0.073 0.117 0.191 0.319
0.0682 0.0815 0.0910 0.0973 0.1029 0.1082 0.0231 0.0390 0.0587 0.0738 0.0849 0.0932 0.0026 0.0046 0.0082 0.0131 0.0212 0.0350
0.972 1.160 1.357 1.495 1.668 1.840 0.435 0.673 0.872 1.018 1.210 1.409 0.052 0.090 0.153 0.233 0.361 0.579
0.0897 0.1052 0.1209 0.1316 0.1447 0.1572 0.0423 0.0638 0.0813 0.0935 0.1093 0.1250 0.0052 0.0091 0.0152 0.0231 0.0353 0.0554
1.686 1.834 1.947 1.992 2.018 2.080 0.625 0.975 1.465 1.718 1.869 1.967 0.059 0.100 0.161 0.233 0.485 0.715
0.1144 0.1232 0.1298 0.1324 0.1339 0.1375 0.0457 0.0695 0.1009 0.1163 0.1253 0.1310 0.0045 0.0076 0.0122 0.0175 0.0358 0.0519
pCO2 = (ptotal − pwater)·CO2 vol %. bu(T) = 0.1 K, u(pCO2) = 0.1 kPa, u(CO2 solubility) = 2 % of (CO2 solubility [molCO2/molamine]). D
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and β include the zero-point energy and thermal corrections dependent on the level of theory.6
Table 4. Experimental and Calculated Values of pKa of Amine at 295 Ka
3. RESULTS AND DISCUSSION Figure 2 shows CO2 solubility measurement results for the three tertiary diamines, and Table 2 shows CO2 solubility data
SMD/B3LYP/6-311+ +G(d,p)
experiment C2-diamine C3-diamine C6-diamine MDEA a
Figure 3. Comparison of CO2 solubility for the three diamines, MDEA and MEA at 313 K, 30 wt % amine: ▲, C2-diamine; ■, C3-diamine; ⧫, C6-diamine; □, MDEA; ◊, MEA with the smoothed line.
pKa1
pKa2
pKa1
pKa2
9.5 10.1
6.1 8.1
9.51 10.4 10.2
6.13 8.28 10.0
10.1 8.5
u(pKa) = 0.1, u(T) = 0.5 K.
Figure 5. Equiliblium constant log10K313 determined with eqs 14 and 15 against pKa.
Figure 6. Reaction enthalpy ΔH determined with eqs 14 and 15 against pKa.
Figure 4. Deviation between experimental and model correlation results for CO2 solubility: blue ◆, C2-diamine; red ■, C3-diamine; green ▲, C6-diamine.
diamine. Sing et al. studied the effect of alkyl chain length of primary diamine on CO2 solubility and showed that CO2 solubility increased according to carbon number between amino groups: C2 < C4 ≤ C3 ≤ C7 < C6.7 They suggested that a formation of a hydrogen bond between the amino groups might play an important role. In addition, this trend implies that the higher CO2 solubility is given by the larger carbon number because of the higher amine basicity.
for the three tertiary diamines. For all diamines, the CO2 solubility increased with decreasing temperature and increasing pressure. Figure 3 shows the comparison of CO2 solubility for the three diamines, MDEA,27 and monoethanolamine (MEA)30 at 313 K. Compared with conventional amines MDEA and MEA, all of the diamines showed high CO2 solubility. In terms of alkyl chain length between the two amines, CO2 solubility increased according to C2-diamine < C3-diamine < C6-
Table 3. K313 and ΔH Value of Amines Determined with eqs 14 and 15
C2-diamine C3-diamine C6-diamine MDEA
K1313/(MPa−1)
ΔH1/(kJ·mol−1)
K2313/(MPa−1)
ΔH2/(kJ·mol−1)
129.7 369.5 431.8 30
−58.7 −68.1 −72.1 −49.1
0.157 9.969 431.8
−40.5 −47.7 −72.1
E
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parameters for each diamine and MDEA.27 The model can represent the CO2 solubility data well. For the equilibrium constant, C2-diamine and C3-diamine have two values, whereas C6-diamine has one value. Table 4 shows the pKa for these tertiary amines. C2-Diamine and C3-diamine have two pKa values, although C6-diamine has one pKa value; this is analogous to the trend for equilibrium constant K. For C2-Diamine, the basicity of one amine will become smaller when the other amine is protonated. Furthermore, C3-diamine exhibits that behavior, but the variation is smaller than that for C2-diamine. For C6-diamine, the basicity for each amine is almost the same, as the alkyl chain length is sufficiently long so that the two amines do not influence each other. As will be discussed, these trends are consistent with predictions of the DFT calculation using the SMD solvation model. Figures 5 and 6 show the equiliblium constant log10K313 determined with eqs 14 and 15 against pKa and the reaction enthalpy ΔH against pKa, respectively. The log10(K313)−pKa relationship is linear; therefore, a higher basicity is correlated with a higher CO2 solubility. Figure 7 shows the reaction enthalpy ΔH determined with the Clausius−Clapeyron equation using the model parameters obtained in this study. The reaction enthalpy increased according to MDEA < C2diamine < C3-diamine < C6-diamine. The reaction enthalpy of C2-diamine and C3-diamine decreased gradually with the CO2 solubility, because of the two pKa values (pKa1 and pKa2 in Table 3). Figure 8 shows the geometrically optimized structure of diprotonated C6-Diamine using the DFT calculation for the aqueous phase. We considered linear conformers, as shown in Figure 8, for all the neutral, protonated, and diprotonated diamines to clearly compare the effects of the alkyl chain length between the two amino groups. All of the energy differences between a particular neutral and protonated diamine or between a protonated and diprotonated diamine were calculated from the DFT optimizations and plotted against the corresponding pKa values, as shown in Figure 9. The calculated energy differences in the aqueous phase showed strong correlation with the experimental pKa values (R2 = 0.99), whereas this correlation was weaker for the vacuum phase (R2 = 0.71). These results demonstrate that the alkyl chain between amino groups enhances amino basicity because of the electrondonating property and indicate that electrostatic repulsion between protonated amino groups in diprotonated diamines can be reduced by the effects of solvation as well as alkyl chain length.
Figure 7. Reaction enthalpy ΔH determined with the Clausius− Clapeyron equation with our model parameters.
Figure 8. Geometrically optimized structure of diprotonated C6diamine by the DFT calculation for the aqueous phase.
4. CONCLUSION The CO2 solubility of three types of tertiary diamine solutions was measured with an equilibrium cell. A longer alkyl chain between the two amines led to higher CO2 solubility, as it depends on amine basicity, which can be explained by the electron-donating property of the alkyl group and the electrostatic effects in aqueous solutions. A CO2 solubility model was constructed by considering the chemical equilibrium constant, and it could represent experimental data well. Equilibrium constant and pKa are linearly correlated.
Figure 9. Energy differences between neutral and protonated diamines or between protonated and diprotonated diamines calculated from the DFT optimizations.
Figure 2 shows the model correlation results, and Figure 4 shows the deviation pot of CO 2 solubility between experimental and correlation data. Table 3 shows model F
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(17) Kumar, G.; Kundu, M. Solubility of CO2 in Aqueous Blends of (N-Methyl-2-ethanolamine + N-Methyl-diethanolamine) and (NMethyl-2-ethanolamine + 2-Amino-2-methyl-1-propanol). J. Chem. Eng. Data 2012, 57, 3203−3209. (18) Kumar, G.; Kundu, M. Vapor−Liquid Equilibrium of CO2 in Aqueous Blends of (N-Ethyl-ethanolamine + N-Methyl-diethanolamine) and (N-Ethyl-ethanolamine + 2-Amino-2-methyl-1-propanol). J. Chem. Eng. Data 2013, 58, 2959−2965. (19) Mandal, B. P.; Kundu, M.; Bandyopadhyay, S. S. Physical Solubility and Diffusivity of N2O and CO2 into Aqueous Solutions of (2-Amino-2-methyl-1-propanol + Monoethanolamine) and (N-Methyldiethanolamine + Monoethanolamine). J. Chem. Eng. Data 2005, 50, 352−358. (20) Mandal, B. P.; Kundu, M.; Padhiyar, N. U.; Bandyopadhyay, S. S. Physical Solubility and Diffusivity of N2O and CO2 into Aqueous Solutions of (2-Amino-2-methyl-1-propanol + Diethanolamine) and (N-Methyldiethanolamine + Diethanolamine). J. Chem. Eng. Data 2004, 49, 264−270. (21) Gabrielsen, J.; Michelsen, M. L.; Stenby, E. H.; Kontogeorgis, G. M. A model for estimating CO2 solubility in aqueous alkanolamines. Ind. Eng. Chem. Res. 2005, 44, 3348−3354. (22) Tobiesen, F. A.; Svendsen, H. F.; Mejdell, T. Modeling of blast furnace CO2 capture using amine absorbents. Ind. Eng. Chem. Res. 2007, 46, 7811−7819. (23) Fouad, W. A.; Berrouk, A. S. Prediction of H2S and CO2 Solubilities in Aqueous Triethanolamine Solutions Using a Simple Model of Kent-Eisenberg Type. Ind. Eng. Chem. Res. 2012, 51, 6591− 6597. (24) Posey, M. L.; Tapperson, K. G.; Rochelle, G. T. A simple model for prediction of acid gas solubilities in alkanolamines. Gas Sep. Purif. 1996, 10, 181−186. (25) Tomizaki, K.; Kanakubo, M.; Nanjo, H.; Shimizu, S.; Onoda, M.; Fujioka, Y. 13)C NMR Studies on the Dissolution Mechanisms of Carbon Dioxide in Amine-Containing Aqueous Solvents at High Pressures toward an Integrated Coal Gasification Combined CycleCarbon Capture and Storage Process. Ind. Eng. Chem. Res. 2010, 49, 1222−1228. (26) Tomizaki, K.; Shimizu, S.; Onoda, M.; Fujioka, Y. An acid dissociation constant (pKa)-based screening of chemical absorbents that preferably capture and release pressurized carbon dioxide for greenhouse gas control. Chem. Lett. 2008, 37, 516−517. (27) Tomizaki, K.; Shimizu, S.; Onoda, M.; Fujioka, Y. Heats of Reaction and Vapor-Liquid Equilibria of Novel Chemical Absorbents for Absorption/Recovery of Pressurized Carbon Dioxide in Integrated Coal Gasification Combined Cycle-Carbon Capture and Storage Process. Ind. Eng. Chem. Res. 2010, 49, 1214−1221. (28) Chowdhury, F. A.; Yamada, H.; Higashii, T.; Goto, K.; Onoda, M. CO2 Capture by Tertiary Amine Absorbents: A Performance Comparison Study. Ind. Eng. Chem. Res. 2013, 52, 8323−8331. (29) Chowdhury, F. A.; Okabe, H.; Shimizu, S.; Onoda, M.; Fujioka, Y. Development of novel tertiary amine absorbents for CO2 capture. Energy Procedia 2009, 1, 1241−1248. (30) Yamada, H.; Chowdhury, F. A.; Goto, K.; Higashii, T. CO2 solubility and species distribution in aqueous solutions of 2(isopropylamino)ethanol and its structural isomers. Int. J. Greenhouse Gas Control 2013, 17, 99−105. (31) Marenich, A. V.; Cramer, C. J.; Truhlar, D. G. Universal Solvation Model Based on Solute Electron Density and on a Continuum Model of the Solvent Defined by the Bulk Dielectric Constant and Atomic Surface Tensions. J. Phys. Chem. B 2009, 113, 6378−6396. (32) Yamada, H.; Matsuzaki, Y.; Higashii, T.; Kazama, S. Density Functional Theory Study on Carbon Dioxide Absorption into Aqueous Solutions of 2-Amino-2-methyl-1-propanol Using a Continuum Solvation Model. J. Phys. Chem. A 2011, 115, 3079−3086.
AUTHOR INFORMATION
Corresponding Author
*Tel. +81-(0)52-789-3622. Fax +81-(0)52-789-3272. E-mail:
[email protected]. Funding
This study was financially supported by a Grant-in-Aid from the Ministry of Economy, Trade and Industry (METI) of Japan. Notes
The authors declare no competing financial interest.
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REFERENCES
(1) Aaron, D.; Tsouris, C. Separation of CO2 from flue gas: A review. Sep. Sci. Technol. 2005, 40, 321−348. (2) D’Alessandro, D. M.; Smit, B.; Long, J. R. Carbon Dioxide Capture: Prospects for New Materials. Angew. Chem., Int. Ed. 2010, 49, 6058−6082. (3) Yang, H. Q.; Xu, Z. H.; Fan, M. H.; Gupta, R.; Slimane, R. B.; Bland, A. E.; Wright, I. Progress in carbon dioxide separation and capture: A review. J. Environ. Sci.-China 2008, 20, 14−27. (4) Rochelle, G. T. Amine Scrubbing for CO2 Capture. Science 2009, 325, 1652−1654. (5) Yamada, H.; Matsuzaki, Y.; Chowdhury, F.; Higashii, T. Computational investigation of carbon dioxide absorption in alkanolamine solutions. J. Mol. Model. 2013, 19, 4147−4153. (6) Yamada, H.; Shimizu, S.; Okabe, H.; Matsuzaki, Y.; Chowdhury, F. A.; Fujioka, Y. Prediction of the Basicity of Aqueous Amine Solutions and the Species Distribution in the Amine-H2O-CO2 System Using the COSMO-RS Method. Ind. Eng. Chem. Res. 2010, 49, 2449− 2455. (7) Singh, P.; Niederer, J. P. M.; Versteeg, G. F. Structure and activity relationships for amine based CO2 absorbents - I. Int. J. Greenh. Gas Con. 2007, 1, 5−10. (8) Singh, P.; Niederer, J. P. M.; Versteeg, G. F. Structure and activity relationships for amine-based CO2 absorbents-II. Chem. Eng. Res. Des. 2009, 87, 135−144. (9) Bottger, A.; Ermatchkov, V.; Maurer, G. Solubility of Carbon Dioxide in Aqueous Solutions of N-Methyldiethanolamine and Piperazine in the High Gas Loading Region. J. Chem. Eng. Data 2009, 54, 1905−1909. (10) Ermatchkov, V.; Kamps, A. P. S.; Maurer, G. Solubility of carbon dioxide in aqueous solutions of N-methyldiethanolamine in the low gas loading region. Ind. Eng. Chem. Res. 2006, 45, 6081−6091. (11) Ermatchkov, V.; Maurer, G. Solubility of carbon dioxide in aqueous solutions of N-methyldiethanolamine and piperazine: Prediction and correlation. Fluid Phase Equilib. 2011, 302, 338−346. (12) Kamps, A. P. S.; Balaban, A.; Jodecke, M.; Kuranov, G.; Smirnova, N. A.; Maurer, G. Solubility of single gases carbon dioxide and hydrogen sulfide in aqueous solutions of N-methyldiethanolamine at temperatures from 313 to 393 K and pressures up to 7.6 MPa: New experimental data and model extension. Ind. Eng. Chem. Res. 2001, 40, 696−706. (13) Kamps, A. P. S.; Xia, J. Z.; Maurer, G. Solubility of CO2 in (H2O + piperazine) and in (H2O + MDEA + piperazine). AIChE J. 2003, 49, 2662−2670. (14) Kuranov, G.; Rumpf, B.; Smirnova, N. A.; Maurer, G. Solubility of single gases carbon dioxide and hydrogen sulfide in aqueous solutions of N-methyldiethanolamine in the temperature range 313− 413 K at pressures up to 5 MPa. Ind. Eng. Chem. Res. 1996, 35, 1959− 1966. (15) Speyer, D.; Ermatchkov, V.; Maurer, G. Solubility of Carbon Dioxide in Aqueous Solutions of N-Methyldiethanolamine and Piperazine in the Low Gas Loading Region. J. Chem. Eng. Data 2010, 55, 283−290. (16) Bougie, F.; Iliuta, M. C. Sterically Hindered Amine-Based Absorbents for the Removal of CO2 from Gas Streams. J. Chem. Eng. Data 2012, 57, 635−669. G
DOI: 10.1021/je500927h J. Chem. Eng. Data XXXX, XXX, XXX−XXX
