Enthalpy of Solution of Carbon Dioxide in Aqueous Solutions of

Article pubs.acs.org/jced

Enthalpy of Solution of Carbon Dioxide in Aqueous Solutions of Triethanolamine at Temperatures of 322.5 K and 372.9 K and Pressures up to 5 MPa Paper presented at the 18th Symposium on Thermophysical Properties, Boulder, CO, June 24 to 29, 2012. Hugues Arcis,*,†,‡ Karine Ballerat-Busserolles,‡ Laurence Rodier,‡ and Jean-Yves Coxam*,‡ †

Department of Chemistry, University of Guelph, Guelph, Ontario, Canada N1G 2W1 Institut Chimie de Clermont Ferrand, Thermodynamique et Interactions Moléculaires, UMR CNRS 6296, Blaise Pascal University, Aubière, France

‡

ABSTRACT: The enthalpies of solution (ΔsolH) of carbon dioxide (CO2) in two aqueous solutions (w = 0.1500 and w = 0.3000) of triethanolamine (TEA) have been measured for the first time at T = 322.5 K and 372.9 K and p up to 5 MPa, by flow calorimetry. Enthalpies of solution of CO2 (ΔsolH) have been obtained as function of loading α (molCO2·molamine−1). Solubility data (s) of the gas into the solution were also determined. Influences of temperature, pressure, and absorbent composition have been discussed. Our experimental enthalpies of solution have been compared with data derived from a thermodynamic model of phase equilibria based on a γ−ϕ approach. The enthalpy of solution of CO2 in aqueous solutions of methyldiethanolamine (MDEA), monoethanolamine (MEA), and diethanolamine (DEA) have been compared to the present results to discuss the effects of the hindrance on the nitrogen nucleus as well as its degree of substitution, on the enthalpy of solution.

1. INTRODUCTION For a couple of decades, aqueous solutions of alkanolamines have been known to be very effective absorbents for the deacidification of natural gas in so-called washing amine processes. Now the aim is to adjust the technology for removing carbon dioxide from industrial postcombustion effluents. The development of capture process is one possible solution found to fight against global warning through the reduction of carbon dioxide emissions. The absorption of carbon dioxide implies simultaneous subsequent chemical reactions and physical dissolution. The acidobasic reactions between the gas and the chemical solvent are reversible, opening a way for a possible gas/solvent separation when working through a cyclic industrial process. The objective of current research on CO2 capture is the diminution of the energetic cost of solvent regeneration, that is, when removing the gas from the absorbent in the stripper unit. To design new absorbents, it is of great importance to understand mechanisms of gas absorption. For this purpose and to get a full picture, we carry out a systematic study on molecules representative of primary, secondary, and terciary amine families. We particularly look at the energies of gas dissolution into aqueous solutions made of common alkanolamines.1−5 This paper is the last one of a series of three that were focused on aqueous solutions of mono-, di-, and triethanolamine, aiming to a better understanding of the impact of the structural properties of the solvent on the energies involved. © 2012 American Chemical Society

The experimental data are used to develop the thermodynamic models representative of {CO2 + water + amine} systems and are required to study gas dissolution and design future CO2 capture processes. A few thermodynamic data on vapor−liquid equilibria for the system {CO2 + TEA + H2O} have been reported in the literature6−10 (see Table 1), and to Table 1. Literature Review of Gas Solubility Data for the System {CO2 + TEA + H2O} molarity reference Jou et al.6 Mason and Dodge7 Shneerson and Leibush8 Lyudkovskaya and Leibush9 Van den Berg10

mol·L

−1

2.0−5.0 0.5−5.0 0.5−3.5 0.5−5.0 0.1−1.0

pCO2

δmax

K

kPa

%

298−398 273−348 323 300−333 276

0.01−6360 1.32−100 0.04−49 255−4124 0.1−9.3

T

2

our knowledge, one study has been realized,11 but no numerical data for experimental enthalpy have been ever published except a single value,12 (989 kJ·kg−1 of CO2) for the enthalpy of reaction. In this paper, we report the very first set of experimental data on the enthalpy of solution for the system {CO2 + TEA + H2O} at Received: July 12, 2012 Accepted: October 30, 2012 Published: November 27, 2012 3587

dx.doi.org/10.1021/je300813c | J. Chem. Eng. Data 2012, 57, 3587−3597

Journal of Chemical & Engineering Data

Article

temperatures higher than 313 K. Measurements were performed at constant temperature (T = 322.5 K and T = 372.9 K) and pressure (p = 0.5 MPa, 1 MPa, 3 MPa, 5 MPa) for two different aqueous solutions (wTEA = 0.1500 ± 0.0001 and wTEA = 0.3000 ± 0.0001), using a calorimetric technique developed in our laboratory.13 Solubilities of CO2 into the aqueous solutions (s) were derived from the enthalpy data. The different contributions to the enthalpy of solution (ΔsolH) of CO2 in the two aqueous solutions of triethanolamine have also been analyzed and discussed, with the help of a thermodynamic model of phase equilibria described previously.5,14 Comparisons among CO2 absorption in mono-, di-, and triethanolamine were made to discuss the effect of the hindrance on the nitrogen nucleus in the case of basic ethanolamines. The effect of the hindrance on the nitrogen nucleus in the case of tertiary amine is based on the comparison between dissolution in triethanolamine (TEA) and methyl-diethanolamine (MDEA) aqueous solutions.

Table 3. Experimental Enthalpies of Solution of CO2 in Aqueous Solutions of TEA (wTEAa = 0.1500) at Tb = 322.50 K α

K

MPa

ρ kg·m

−3

Ta

pb

ρ

K

MPa

kg·m−3

wTEAc 298.31 298.31 298.32 298.31 298.32 298.32 a

= 0.1500 0.206 1019.8 0.505 1020.0 1.008 1020.2 2.002 1020.6 3.001 1021.0 5.007 1021.8

c

± ± ± ± ± ±

0.1 0.1 0.1 0.1 0.1 0.1

298.32 298.31 298.31 298.31 298.32 298.31

wTEA = 0.3000 0.203 1035.5 0.508 1035.6 1.002 1035.8 2.007 1036.2 3.002 1036.5 5.005 1037.3

± ± ± ± ± ±

δΔsolH

−ΔsolH

−1

−1

kJ·mol kJ·mol kJ·mol molCO2·molTEA−1 molCO2·molTEA−1 of TEA of TEA of CO2

Table 2. Experimental Density of Aqueous TEA Solutions (w = 0.1500 and 0.3000) Used for Calculating the Molar Flow Rates of the Amine Solutions pb

−ΔsolH −1

2. EXPERIMENTAL SECTION The experimental setup has already been described in previous publications when studying the dissolution of CO2 in aqueous solutions of alkanolamines,1−5 The enthalpy of solution of CO2 in the aqueous solutions of amine was measured using a custom-made flow-mixing cell adapted to a Setaram C-80 heat conduction differential calorimeter.1−5,13 Experiments were carried out at constant temperature and pressure as a function of the gas loadings, α (molCO2·molamine−1). Uncertainties were estimated following previous practices1−5 and are reported in Tables 3 to 6 for each experimental data point. Densities of aqueous TEA solutions at T = 298 K and different pressures required in data analysis, were obtained using an Anton Paar densimeter DMA 512 (P model);15 values are reported in Table 2.

Ta

δα

0.1 0.1 0.1 0.1 0.1 0.1

u(T) = 0.01 K. bu(p) = 0.002 MPa. cu(w) = 0.0001.

The densities of CO2 were calculated from ALLPROPS software.16 Triethanolamine was purchased from Fluka organics with a purity of > 99 % and was used without further purification. Carbon dioxide (purity of 99.998 %) was obtained from Saga. Solutions were prepared by dissolution in distilled and degassed water (resistivity 18.2 MΩ·cm) with an uncertainty in the mass fraction estimated to less than ± 10−4.1−5

3. EXPERIMENTAL RESULTS Experimental enthalpies of solution of CO2 in aqueous TEA solutions are reported as a function of the loading (α = molCO2·molTEA−1) in Tables 3 to 6, and the results obtained at T = 372.9 K for wTEA = 0.300 presented as example in Figure 1. 3588

0.286 0.334 0.343 0.389 0.445 0.491 0.546 0.601 0.667 0.710 0.764 0.850 0.963 1.019 1.076 1.133 1.268 1.268 1.384 1.586 1.845

0.005 0.006 0.006 0.007 0.008 0.009 0.010 0.011 0.012 0.013 0.014 0.016 0.018 0.019 0.020 0.021 0.023 0.023 0.025 0.029 0.033

0.182 0.228 0.304 0.364 0.451 0.511 0.602 0.694 0.780 0.876 0.974 1.062 1.067 1.154 1.242 1.343 1.454 1.580 1.709

0.002 0.003 0.004 0.004 0.005 0.006 0.007 0.009 0.010 0.011 0.012 0.013 0.012 0.014 0.015 0.016 0.017 0.019 0.020

0.490 0.567 0.648 0.783 0.826 0.977 1.012 1.046 1.072 1.113 1.121 1.147 1.159 1.186 1.272

0.003 0.003 0.003 0.004 0.004 0.005 0.005 0.006 0.006 0.006 0.006 0.006 0.006 0.006 0.007

pc = 0.54 MPa 14.2 16.1 16.1 18.7 21.7 23.3 25.4 27.7 29.9 32.5 35.1 37.5 39.7 39.9 40.1 39.2 39.2 39.4 40.2 39.2 39.5 pc = 1.06 MPa 8.5 10.6 14.2 16.6 21.0 24.2 28.7 32.5 36.1 40.3 44.1 47.3 47.6 46.2 46.9 46.8 47.1 47.1 47.3 pc = 5.31 MPa 22.8 24.6 29.6 34.3 37.4 44.6 42.8 45.9 47.3 47.9 49.6 47.9 51.2 50.0 51.2

δΔsolH kJ·mol−1 of CO2

0.1 0.3 0.3 0.5 0.2 0.5 0.5 0.4 0.4 0.4 0.6 0.5 0.5 0.5 0.6 0.5 0.6 0.5 0.9 0.5 0.5

49.8 48.1 46.9 48.1 48.7 47.4 46.5 46.1 44.8 45.8 46.0 44.1 41.2 39.1 37.3 34.6 30.9 31.1 29.0 24.7 21.4

0.3 1.0 0.9 1.3 0.5 1.1 0.8 0.7 0.5 0.6 0.8 0.6 0.6 0.5 0.5 0.4 0.4 0.4 0.7 0.3 0.3

0.1 0.2 0.3 0.5 0.3 0.6 0.5 0.8 0.5 0.6 0.7 0.5 0.5 0.7 0.6 0.6 0.6 0.6 0.6

46.6 46.4 46.8 45.5 46.6 47.3 47.7 46.8 46.3 46.1 45.3 44.6 44.6 40.1 37.8 34.9 32.4 29.8 27.7

0.7 0.9 0.9 1.3 0.7 1.2 0.9 1.2 0.6 0.7 0.7 0.5 0.4 0.6 0.5 0.5 0.4 0.3 0.3

0.2 0.4 0.4 0.3 0.2 1.2 1.3 1.0 0.3 0.7 0.5 0.8 1.4 0.8 1.2

46.5 43.4 45.7 43.8 45.2 45.6 42.3 43.8 44.1 43.1 44.2 41.8 44.2 42.1 40.3

0.4 0.8 0.6 0.4 0.3 1.3 1.3 1.0 0.3 0.7 0.4 0.7 1.2 0.7 0.9



Article

Table 3. continued α

δα

−ΔsolH

δΔsolH

−1

−ΔsolH

−1

−1

kJ·mol kJ·mol kJ·mol molCO2·molTEA−1 molCO2·molTEA−1 of TEA of TEA of CO2 1.336 1.399 1.558 1.766 2.026 2.197 2.373 a

0.007 0.007 0.008 0.009 0.011 0.012 0.012

pc = 5.31 MPa 55.1 55.0 55.9 55.8 55.7 56.0 55.4

0.4 0.9 0.4 0.4 0.4 0.5 0.8

41.3 38.6 35.9 31.6 27.5 25.5 23.3

Table 4. Experimental Enthalpies of Solution of CO2 in Aqueous Solutions of TEA (wTEAa = 0.3000) at Tb = 322.50 K

δΔsolH

α

kJ·mol−1 of CO2 0.3 0.6 0.3 0.3 0.2 0.2 0.3

We note that our results are consistent with data of Kwak et al.11 In all cases, the enthalpy of solution was observed to be constant within the experimental uncertainty up to the limit of solubility (s/molCO2·molTEA−1) where s corresponds to the loading when the solution is saturated in gas. This trend agrees with literature description of chemical absorption of CO2 in aqueous solutions of tertiary amines.17,18 Barth et al.19 reported the following mechanism to explain the dissolution of CO2 into aqueous solutions of amine lacking a hydrogen on the amino group: (1)

(C2H4OH)3 N+COO− + H 2O ⇄ (C2H4OH)3 NH+ + HCO−3

(2)

The overall reaction can be then considered as: (C2H4OH)3 N + CO2 + H 2O ⇆ (C2H4OH)3 NH+ + HCO

−ΔsolH

molCO2·molTEA−1 molCO2·molTEA−1

u(w) = 0.0001. bu(T) = 0.03 K. cu(p) = 0.01 MPa.

(C2H4OH)3 N + CO2 ⇄ (C2H4OH)3 N+COO−

δα

(3)

The fact that the absorption of CO2 is governed by the formation of bicarbonate would explain why the enthalpy remains constant up to the solubility limit, in opposition to what was observed for primary,1,4 and secondary amine,5 where the absorption of the gas into the solution was governed by a competition between the formation of bicarbonate and carbamate. Moreover, despite the stoichiometry limit of 1 mol of CO2 per mole of amine, the solubility of the gas can be extended to loading higher than 1 when pressure increases if we consider physical absorption. 3.1. Solubility Data. Following previous practice,1−5 the solubility limits of the gas into the absorbent at different conditions of temperature and pressure were obtained by determining the split point between the unsaturated (linear increase) and saturated domains (plateau) when the enthalpy is expressed in kJ per mol of amine. As an example we plotted in Figure 1 our enthalpy data for the absorption of CO2 at T = 372.9 K and for a solution wTEA = 0.3000. Numerical values for solubility limit were graphically determined from Figure 1 and are reported in Table 7. Following previous studies,1−5 our solubility limits were assumed to be determined to within ± 5 %. Our results are shown in Figure 2, together with literature values from Jou et al.6 Our solubility limits derived from experimental enthalpy data and theirs obtained by specific technics devised for VLE property determination are in excellent agreement except at 373 K and 3 MPa, where we note a difference of ∼ 10 %. We also compared in Figure 3 our solubility limits of CO2 in our aqueous solutions of TEA with solubility limits of CO2 in aqueous solutions 3589

0.328 0.437 0.492 0.546 0.601 0.655 0.710 0.765 0.819 0.874 0.983 1.092

0.006 0.008 0.009 0.010 0.011 0.012 0.013 0.014 0.015 0.016 0.018 0.020

0.201 0.242 0.260 0.322 0.421 0.526 0.564 0.625 0.645 0.729 0.767 0.833 0.946 1.052 1.244 1.401 1.503 1.615 1.902

0.003 0.003 0.003 0.004 0.005 0.007 0.007 0.008 0.008 0.009 0.010 0.011 0.012 0.014 0.013 0.015 0.018 0.016 0.018

0.499 0.539 0.592 0.652 0.673 0.802 0.841 0.863 0.891 0.914 0.925 0.941 0.960 0.972 1.020 1.031 1.077 1.100 1.176 1.266 1.296 1.379 1.509

0.003 0.003 0.003 0.003 0.004 0.004 0.004 0.005 0.005 0.005 0.005 0.005 0.005 0.005 0.005 0.005 0.006 0.006 0.006 0.007 0.007 0.007 0.008

δΔsolH

−ΔsolH

δΔsolH

kJ·mol−1 kJ·mol−1 kJ·mol−1 kJ·mol−1 of TEA of TEA of CO2 of CO2

pc = 0.52 MPa 16.7 22.1 25.4 28.0 29.6 28.7 29.8 28.6 28.6 29.8 29.5 29.8 pc = 1.03 MPa 10.4 11.9 13.2 15.8 21.9 26.1 27.7 30.8 31.7 35.5 36.7 39.7 41.6 42.2 42.3 42.2 42.2 42.9 42.3 pc = 5.36 MPa 24.6 27.2 29.9 32.8 33.7 39.4 41.1 43.3 43.1 45.1 44.5 46.2 44.5 47.4 47.3 47.2 49.6 49.4 50.6 51.7 51.0 51.6 51.5

0.1 0.2 0.2 0.2 0.3 0.3 0.3 0.3 0.3 0.3 0.3 0.4

50.8 50.5 51.6 51.3 49.2 43.8 41.9 37.4 34.9 34.1 30.0 27.3

0.4 0.5 0.5 0.4 0.5 0.5 0.4 0.4 0.4 0.4 0.3 0.4

2.0 0.2 0.2 0.4 0.3 0.4 0.4 0.4 0.4 0.4 0.9 0.4 0.4 0.6 0.7 0.4 0.5 0.4 0.4

51.5 49.2 50.6 49.0 51.9 49.6 49.2 49.3 49.2 48.6 47.8 47.7 43.9 40.1 34.0 30.1 28.1 26.6 22.3

0.8 0.8 0.8 1.2 0.8 0.8 0.8 0.6 0.6 0.5 1.1 0.5 0.5 0.6 0.5 0.3 0.4 0.3 0.2

0.3 0.5 0.6 0.3 0.3 2.5 0.6 0.3 0.6 4.1 0.9 0.4 0.6 1.1 0.4 0.3 0.7 0.4 0.2 0.2 0.3 0.2 0.5

49.3 50.4 50.6 50.3 50.1 49.1 48.9 50.1 48.3 49.3 48.1 49.1 46.4 48.8 46.3 45.8 46.1 44.9 43.0 40.9 39.4 37.4 34.1

0.7 1.0 1.0 0.4 0.4 3.1 0.7 0.3 0.7 4.5 1.0 0.4 0.6 1.2 0.4 0.3 0.7 0.3 0.2 0.2 0.2 0.2 0.3



Article


δα


δΔsolH

−ΔsolH

δΔsolH

α


pc = 5.36 MPa 0.009 50.5 0.009 51.4

1.627 1.644 a

−ΔsolH


0.3 0.4

31.1 31.2

=

∫p

1

⎡ • ⎞ ⎤ ⎛ ∂V CO ⎢ • 2 ⎜ ⎟⎥ ⎢V CO2 − T ⎜⎝ ∂T ⎟⎠ ⎥dp ⎣ p⎦

0.2 0.2

(4)

The residual enthalpy of CO2 in the condition of infinite dilution in the aqueous phase (eq 5) was obtained using the aqueous molar volume of CO2 at infinite dilution reported by Silkenbäumer et al.19 ∞ ΔHres

=

∫p

1

p2

⎡ ⎛ ∞ ⎞⎤ ⎢V ∞ − T ⎜ ∂V CO2 ⎟ ⎥dp ⎜ ⎟ ⎢ CO2 ⎝ ∂T ⎠ p⎥⎦ ⎣

(5)

At both temperatures, a linear diminution of the exothermic effect with pressure was obtained. This diminution corresponds to the residual enthalpies (eq 6). • ∞ Δsol H ∞(T , p1 ) = Δsol H ∞(T , p2 ) + ΔHres − ΔHres

δΔsolH

−ΔsolH

−1

−1


of MDEA, obtained in the same experimental conditions.2,3 The solubility of CO2 in the aqueous solutions of TEA and MDEA were found to follow same trends: decreasing as temperature increases; this can be related to the diminution of amine basicity. Looking at the effect of the composition of the aqueous solution, we observed a decrease of the CO2 solubility (expressed in mol of CO2 per mol of amine), resulting in the diminution of the number of solvent molecules available for solvating the gas because of the increase of ions in solution. We note that the loadings at solubility limit, s, are higher in aqueous solutions of MDEA than of TEA. This has to be expected when considering the substitution of a methyl group by an ethanol group when comparing MDEA to TEA, which leads to a decrease of TEA pKa value compared to the one for MDEA. The difference in solubility was also found to increase with temperature and to decrease with CO2 partial pressure at equilibrium. 3.2. Enthalpy Data at Low Loadings. At low loadings (Figure 4), the enthalpy of solution per mole of CO2 was found to be constant, within the experimental uncertainty as observed in previous literature study.1−5 Average enthalpy values (ΔsolHav/kJ·molCO2−1) for α < s are presented in Table 8 and Figure 4. A small rise of the exothermic effect is observed as the temperature goes up, and a minor diminution when the pressure is increased. In the first case, the temperature effect on the enthalpy was attributed to the change of amine basicity with temperature and will be discussed with the help of a thermodynamic model in Section 5. In the second case, the decrease of the exothermic effect with pressure could be explained by the pressure dependence on the residual enthalpies of CO2 in the gas and liquid phases2 (Figure 5). The residual enthalpy of CO2 in the gas phase (eq 4) was calculated using the ALLPROPS software.16 p2

−ΔsolH −1

u(w) = 0.0001. bu(T) = 0.03 K. cu(p) = 0.01 MPa.

• ΔHres

δα

(6)

Regarding the dependence of the exothermic effect with with the TEA composition, a slight increase was observed (Figure 4). 3590

0.054 0.106 0.106 0.114 0.153 0.170 0.176 0.204 0.209 0.223 0.250 0.259 0.264 0.284 0.309 0.312 0.341 0.369 0.386 0.422 0.454 0.559 0.671 0.783 0.879

0.001 0.002 0.002 0.002 0.003 0.003 0.003 0.004 0.004 0.004 0.005 0.005 0.005 0.005 0.006 0.006 0.006 0.007 0.007 0.008 0.008 0.010 0.012 0.014 0.017

0.057 0.112 0.169 0.169 0.171 0.213 0.223 0.225 0.225 0.241 0.241 0.281 0.281 0.300 0.320 0.335 0.365 0.426 0.436 0.446 0.446 0.490 0.528 0.538 0.542 0.549 0.634 0.741 0.845 1.056 1.254

0.001 0.001 0.002 0.002 0.002 0.003 0.003 0.003 0.003 0.003 0.003 0.003 0.003 0.004 0.004 0.004 0.005 0.005 0.005 0.005 0.005 0.006 0.007 0.007 0.007 0.007 0.008 0.009 0.011 0.013 0.016

pc = 0.55 MPa 1.8 4.8 5.5 5.7 7.7 8.6 8.8 10.3 10.4 10.8 12.8 12.8 13.1 14.3 13.9 14.6 14.0 13.5 13.8 13.5 13.2 12.5 11.1 10.0 9.2 pc = 1.07 MPa 2.8 5.5 8.6 8.6 8.5 10.5 10.9 11.5 11.1 11.9 12.2 14.2 14.5 14.3 15.9 16.5 17.7 20.9 20.8 20.5 21.7 23.9 24.0 23.5 23.4 24.3 23.8 23.3 23.5 22.2 21.4


0.1 0.8 0.1 0.2 0.1 0.4 0.2 0.3 0.1 0.2 0.2 0.3 0.2 0.1 0.2 0.1 0.2 0.2 0.6 0.5 0.1 0.2 0.1 0.2 0.2

50.2 50.2 51.5 50.3 50.4 50.3 49.9 50.7 49.7 48.3 51.3 49.3 49.6 50.2 44.9 46.9 41.0 36.7 35.6 32.1 29.0 22.4 16.5 12.8 10.4

1.8 7.3 1.4 1.7 0.5 2.3 1.1 1.5 0.5 0.7 0.7 1.0 0.9 0.5 0.5 0.5 0.6 0.5 1.5 1.2 0.3 0.4 0.2 0.2 0.2

0.0 0.1 0.1 0.1 0.3 0.3 0.1 0.3 0.2 0.4 0.4 0.2 0.4 0.3 0.6 0.7 0.3 0.3 0.3 0.3 0.3 0.6 0.3 0.4 0.2 0.3 0.3 0.5 0.3 0.4 0.8

48.8 48.7 51.3 50.8 49.9 49.1 48.9 51.0 49.5 49.3 50.6 50.4 51.5 47.6 49.8 49.3 48.5 49.0 47.8 46.1 48.6 48.8 45.3 43.7 43.1 44.3 37.5 31.5 27.8 21.0 17.1

0.9 0.9 0.6 0.9 1.5 1.4 0.3 1.4 1.1 1.6 1.6 0.7 1.4 0.9 1.8 2.0 0.9 0.8 0.7 0.6 0.6 1.2 0.5 0.8 0.3 0.6 0.5 0.6 0.4 0.4 0.6



Article


δα

−ΔsolH


a

2.477

0.030

0.251 0.312 0.370 0.370 0.465 0.544 0.608 0.664 0.667 0.704 0.751 0.755 0.755 0.781 0.834 0.885 0.932 0.959 1.025 1.048 1.091 1.103 1.143 1.170 1.367 1.436 1.479 1.662 1.783 1.884 2.248 2.960 3.632

0.002 0.002 0.003 0.003 0.004 0.004 0.005 0.005 0.005 0.006 0.006 0.006 0.006 0.006 0.007 0.007 0.007 0.008 0.008 0.008 0.009 0.009 0.010 0.010 0.011 0.012 0.012 0.013 0.015 0.016 0.018 0.024 0.030

δΔsolH

−ΔsolH


δΔsolH

α


pc = 1.07 MPa 16.6 pc = 3.20 MPa 11.8 14.8 17.7 17.6 21.8 26.5 28.0 30.9 31.4 33.4 35.9 35.6 36.0 37.4 38.7 42.9 43.2 44.9 46.5 48.1 47.6 47.5 47.8 47.0 47.2 46.5 46.8 46.6 46.3 46.0 45.3 44.6 43.3

0.3

6.7

0.1

0.4 0.6 0.1 0.4 0.2 0.4 0.2 0.3 0.4 0.1 0.4 1.5 0.2 1.3 0.3 0.3 1.4 0.2 0.5 0.3 0.3 0.4 0.3 0.3 0.3 0.3 0.3 0.3 0.3 0.3 0.3 0.3 0.3

47.2 47.4 47.9 47.5 46.9 48.6 46.0 46.5 47.1 47.4 47.9 47.1 47.7 47.9 46.3 48.5 46.4 46.9 45.3 45.9 43.6 43.1 41.8 40.2 34.5 32.4 31.7 28.1 25.9 24.4 20.1 15.1 11.9

1.5 1.8 0.3 1.1 0.3 0.8 0.4 0.4 0.6 0.2 0.5 2.0 0.3 1.6 0.3 0.3 1.5 0.3 0.5 0.3 0.2 0.4 0.2 0.2 0.2 0.2 0.2 0.2 0.1 0.1 0.1 0.1 0.1

CO2 (aq) + H 2O ⇄

(7)

+H

+

HCO3− ⇄ CO32 − + H+

(8) (9)

(C2H4OH)3 N + H+ ⇄ (C2H4OH)3 NH+

(10)

CO2 L ⇆ CO2 V

(11)

H 2OL ⇆ H 2OV

(12)

(C2H4OH)3 NL ⇆ (C2H4OH)3 NV

(13)

δΔsolH −1

−ΔsolH −1


4. PREDICTION OF THE ENTHALPY OF SOLUTION OF CO2 INTO AQUEOUS TEA SOLUTIONS The thermodynamic model used to predict the enthalpy of solution of CO2 into our aqueous solutions of TEA has been described previously.5,14 Briefly, the modeling of phase equilibria in the system {TEA + CO2 + H2O} was based on the following system of equations: HCO3−

−ΔsolH −1

u(w) = 0.0001. bu(T) = 0.03 K. cu(p) = 0.01 MPa.

H 2O ⇄ H+ + OH−

δα

3591

0.083 0.091 0.104 0.112 0.117 0.133 0.133 0.134 0.140 0.160 0.176 0.183 0.196 0.198 0.206 0.224 0.224 0.229 0.229 0.236 0.253 0.266 0.280 0.308 0.343 0.392 0.449 0.506 0.574 0.608

0.002 0.002 0.002 0.002 0.002 0.003 0.003 0.002 0.003 0.003 0.003 0.003 0.004 0.004 0.004 0.004 0.004 0.004 0.004 0.004 0.005 0.005 0.005 0.006 0.006 0.007 0.008 0.009 0.010 0.011

0.102 0.152 0.205 0.250 0.300 0.319 0.360 0.420 0.525 0.630 0.720 0.814 0.916 1.009 2.095 3.074 3.941 4.960 6.023 6.814 7.647 8.425 11.50 13.67 17.38 22.76

0.001 0.002 0.003 0.003 0.004 0.004 0.005 0.005 0.007 0.008 0.009 0.010 0.012 0.013 0.026 0.037 0.048 0.058 0.069 0.080 0.090 0.101 0.133 0.160 0.212 0.266

pc = 0.54 MPa 4.4 4.7 5.6 5.9 6.1 7.2 7.2 7.0 7.5 8.4 9.4 9.5 10.1 10.4 10.7 11.4 11.3 11.1 11.3 11.1 11.0 10.6 10.9 10.7 10.4 9.8 9.4 8.8 8.3 7.6 pc = 1.06 MPa 5.4 8.0 11.0 12.7 15.2 15.3 15.3 15.3 15.0 14.7 13.8 13.4 13.0 12.6 9.3 5.8 2.4 −1.3 −4.3 −8.0 −11.6 −14.5 −24.2 −33.4 −49.1 −68.1


0.1 0.1 0.1 0.1 0.1 0.0 0.0 0.1 0.5 0.1 0.3 0.2 0.2 0.1 0.1 0.1 0.1 0.2 0.1 0.3 0.2 0.1 0.1 0.1 0.2 0.1 0.1 0.1 0.3 0.1

52.7 51.5 54.0 52.7 52.3 53.8 53.8 51.9 53.4 52.2 53.5 51.9 51.6 52.4 51.9 51.0 50.4 48.6 49.4 47.1 43.6 39.9 38.8 34.7 30.3 25.1 20.9 17.4 14.4 12.6

0.6 0.6 0.5 1.3 1.2 0.3 0.3 0.6 3.7 0.4 1.6 1.1 0.8 0.6 0.5 0.4 0.3 1.1 0.5 1.2 0.6 0.3 0.4 0.4 0.6 0.3 0.3 0.2 0.5 0.2

0.2 0.1 0.9 0.1 0.2 0.3 0.2 0.3 0.2 0.3 0.2 0.2 0.2 0.4 0.3 0.4 0.4 0.4 0.6 0.8 0.8 0.9 1.0 0.4 1.1 0.9

53.4 52.5 53.4 50.7 50.8 47.9 42.6 36.5 28.5 23.3 19.2 16.4 14.2 12.5 4.4 1.9 0.6 −0.3 −0.7 −1.2 −1.5 −1.7 −2.1 −2.4 −2.8 −3.0

1.5 1.0 4.2 0.6 0.6 0.9 0.6 0.7 0.3 0.5 0.3 0.2 0.2 0.4 0.1 0.1 0.1 0.1 0.1 0.1 0.1 0.1 0.1 0.0 0.1 0.0



Article


δα

−ΔsolH


δΔsolH

−ΔsolH

Table 7. Experimental Values for the Solubility of CO2 in Aqueous TEA Solutions (wTEA = 0.1500 and 0.3000) at T = 322.5 K and 372.9 K

δΔsolH


pa

sb

MPa

molCO2·molamine−1

c

0.210 0.210 0.281 0.349 0.449 0.516 0.579 0.644 0.664 0.680 0.707 0.745 0.771 0.803 0.872 1.036 1.233 1.376 1.605 1.809 a

0.002 0.002 0.002 0.003 0.004 0.004 0.005 0.005 0.005 0.006 0.006 0.006 0.006 0.007 0.007 0.009 0.010 0.012 0.013 0.015

p = 3.11 MPa 10.5 10.7 14.7 18.0 23.0 26.5 29.4 32.1 33.3 33.8 34.1 34.9 35.5 35.3 35.1 34.8 34.3 34.1 33.8 33.7

0.1 0.3 0.1 0.1 0.5 0.4 0.5 0.2 0.8 0.8 0.2 0.3 0.2 0.3 0.2 0.2 0.3 0.4 0.2 0.3

49.9 50.7 52.1 51.7 51.3 51.4 50.7 49.9 50.2 49.7 48.2 46.8 46.0 44.0 40.3 33.6 27.8 24.8 21.1 18.6

0.7 1.4 0.5 0.4 1.0 0.8 0.8 0.3 1.2 1.1 0.2 0.4 0.3 0.3 0.3 0.2 0.2 0.3 0.1 0.1

w = 0.1500

a

0.54 1.06 5.31

0.963 1.062 1.336

0.55 1.07 3.20

0.284 0.490 1.025

δs

pa

sb

MPa

molCO2·molamine−1

c

w = 0.3000 T = 322.5 Kd 0.05 0.52 0.05 1.03 0.07 5.36 T = 372.9 Kd 0.02 0.54 0.02 1.06 0.05 3.11

δs

c

0.601 0.833 1.266

0.03 0.04 0.06

0.224 0.300 0.664

0.01 0.02 0.03

u(p) = 0.01 MPa. bur(s) = 0.05. cu(w) = 0.0001. du(T) = 0.03 K.

u(w) = 0.0001. bu(T) = 0.03 K. cu(p) = 0.01 MPa.

Figure 2. Solubility of CO2 in aqueous TEA solutions vs total pressure: ◆, this work, wTEA = 0.3000, T = 322.5 K; ●, this work, wTEA = 0.3000, T = 372.9 K; ◇, Jou et al.,6 wTEA = 0.300, T = 323.15 K; ○, Jou et al.,6 wTEA = 0.300, T = 373.15 K.

The model was adjusted with both amine protonation constants. The speciations (Figure 6) derived from the model at temperatures (T = 322.5 and T = 372.9 K) and compositions (wTEA = 0.1500 and wTEA = 0.3000) did not depend significantly with the choice of the protonation constant. However, the enthalpy of solution of CO2 in aqueous TEA solutions were derived from the model following the procedure of Arcis et al.5,14 using each equilibrium constant for the amine protonation. Results for the absorption of CO2 at T = 322.5 K and for wTEA = 0.1500 are plotted as an example in Figure 7. At each temperature and composition, no noticeable difference (Figure 7) was observed between the two models; they followed the same trend and agreed to each other within 2 %. As the loadings increased to the value α = 0.5, the models were found to slightly underestimate the experimental values with a difference that could reach 10 % for α = 1.0.

Figure 1. Enthalpy of solution (−(ΔsolH/(kJ·mol−1 of TEA)) versus CO2 loading for aqueous solution of TEA at T = 372.9 K and wTEA = 0.3000: □, 0.5 MPa; ◇, 1.0 MPa; ○, 3.0 MPa.

The amine (TEA) was assumed to be nonvolatile (y(C2H4OH)3N = 0), resulting in the omission of eq 13. A γ−ϕ approach5,14 was used to describe the phase equilibria in the {CO2 + H2O + TEA} system; the vapor phase was represented by a truncated virial equation and the liquid phase by a Pitzer model, taking into account only binary interactions. Binary interaction parameters βi,j0 and β1i,j were assumed to be the same as for the system {CO2 + H2O + MDEA}.14 We only used one of the sets of interaction parameters reported in the paper:14 the ones obtained when regressing the {CO2 + H2O + MDEA} VLE data using Oscarson et al.,20 protonation constant. They were chosen because of the extensive experimental range (298 K to 598 K). In case of {CO2 + H2O + TEA} system, two different formulations were tested for the amine protonation equilibrium constant (eq 10) (Bates and Allen21 and Simond et al.,22). Both authors derived their values from potentiometric data.

5. DISCUSSION 5.1. Influence of the Equilibrium Constant Used for Amine Protonation on the Prediction of the Enthalpy of Solution. Two equilibrium constants for the amine protonation (Bates and Allen21 and Simond et al.,22) were used to calculate the enthalpy of solution of CO2 in aqueous TEA solutions up to the solubility limit of CO2. The resulting predicted enthalpies of solution were compared with our experimental data in Figure 7. We note that the choice of the formulation used in the model for the protonation constant did not have a 3592



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Table 8. Average Values for the Enthalpies of Solution at Low Loadings in Aqueous TEA Solutions (wTEA = 0.1500 and 0.3000) at T = 322.5 K and 372.9 K pa

−ΔsolHav

δΔsolHav

pa

−ΔsolHav

δΔsolHav

MPa

kJ·molCO2−1

kJ·molCO2−1

MPa

kJ·molCO2−1

kJ·molCO2−1

b

T = 322.5 K

a

0.54 1.06 5.31

wTEAc = 0.1500 47.1 46.6 45.0

0.55 1.07 3.20

wTEAc = 0.1500 50.1 49.7 47.3

1.4 0.52 1.4 1.03 1.4 5.36 Tb = 372.9 K 1.5 1.5 1.4

0.54 1.06 3.11

wTEAc = 0.3000 50.7 49.9 49.5

1.5 1.5 1.5

wTEAc = 0.3000 52.4 52.2 50.7

1.6 1.6 1.5

u(p) = 0.01 MPa. bu(T) = 0.03 K. cu(w) = 0.0001.

Figure 5. Thermodynamic cycle that linked the enthalpy of solution at the standard pressure p0 to the one at a pressure p is represented at a given temperature T.

Figure 3. Solubility of CO2 in aqueous TEA solutions vs total pressure. (a) wAmine = 0.1500: △, this work, T = 322.5 K and wTEA = 0.1500; □, Arcis et al.,2 T = 322.5 K and wMDEA = 0.1500; ○, this work, T = 372.9 K and wTEA = 0.1500; ◇, Arcis et al.,3 T = 372.9 K and wMDEA = 0.1500. (b) wAmine = 0.3000: ▲, this work, T = 322.5 K and wTEA = 0.3000; ■, Arcis et al.,2 T = 322.5 K and wMDEA = 0.3000; ●, this work, T = 372.9 K and wTEA = 0.3000; ◆, Arcis et al.,3 T = 372.9 K and wMDEA = 0.3000.

general trend of the experimental enthalpy curves, with a plateau less pronounced at low loadings. For wTEA = 0.1500, the thermodynamic model was able to predict the enthalpy of solution within less than 5 % at both temperatures (T = 322.5 K and 372.9 K) and within 20 % for the concentrated solutions of TEA (wTEA = 0.3000). To analyze the cause of this behavior, the different enthalpy contributions of every single chemical/physical reactions involved in the dissolution of CO2 in aqueous TEA solutions were detailed and are shown in Figure 8. At both temperatures (T = 322.5 K and T = 372.9 K) the amine protonation (eq 10) clearly appeared to provide the most important contribution with more than 80 % of the total energetic effect and the smallest being due to the formation of the bicarbonate (eq 8). The enthalpy associated to the physical dissolution of CO2 (eq 11) remained constant at ∼ −15 kJ·mol−1 (T = 322.5 K) and ∼ −2.5 kJ·mol−1 (T = 372.9 K). At T = 322.5 K (Figure 8), the enthalpy released by the formation of the carbonate (eq 9) was found to be rather small and endothermic, oscillating respectively between 3 kJ·mol−1 and 4 kJ·mol−1. On the other hand, at T = 372.9 K, the formation of the carbonate became exothermic and counterbalanced the fact that the enthalpy associated to the physical dissolution of CO2 was lower than at 322.5 K. Finally, when looking at the speciation diagram (Figure 6), we noted that, from T = 322.5 K to T = 372.9 K, the quantity of protonated amine decreased by ∼10 %, resulting in a decrease of the enthalpy of solution as observed experimentally. 5.2. Effect of the Hindrance on the Nitrogen Nucleus in the Case of Tertiary Amines for CO2 Absorption: Comparison between TEA and MDEA. Methyldiethanolamine (MDEA) and the triethanolamine (TEA) are tertiary amines

Figure 4. Average enthalpy of solution (−(ΔsolHav/(kJ·mol−1 of CO2)) versus pressure for aqueous solutions of TEA: ◇, this work, T = 322.5 K and wTEA = 0.1500; ◆, this work, T = 322.5 K and wTEA = 0.3000; ○, this work, T = 372.9 K and wTEA = 0.1500; ●, this work, T = 372.9 K and wTEA = 0.3000.

significant influence on the prediction of the enthalpy of solution. In both cases, the same trend was observed, and the predictions agreed to each other within less than 3 %, which is below the experimental uncertainty. The model reproduced the 3593



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Figure 6. Speciation at equilibrium for the system {CO2 + TEA + H2O} calculated from our thermodymanic model. ···, (C2H4OH)3N; , (C2H4OH)3NH+; −−, HCO3−; ---, CO32‑; −·−, CO2. (a) T = 322.5 K and wTEA = 0.1500; (b) T = 322.5 K and wTEA = 0.3000; (c) T = 372.9 K and wTEA = 0.1500; (d) T = 372.9 K and wTEA = 0.3000.

dissolution of carbon dioxide. For pressures higher than 1 MPa we can observe at T = 322.5 K a better gas solubility in TEA solutions (Figure 3). In the case of tertiary amines, the enthalpy of solution per mole of CO2 depends mainly on the enthalpy related to the amine protonation,14 and the enthalpy value for the amine protonation is itself dependent on the trend of the derivative of the equilibrium constant of protonation (Ka) with respect to the temperature (standard enthalpy of reaction), following van’t Hoff relationships. The Ka constants for TEA and MDEA were determined as a function of temperature in a previous work.23 The results show that the standard enthalpy of reaction for the amine protonation is slightly lower in the case of TEA compared to MDEA. This could contribute to a lower enthalpy of solution of carbon dioxide in TEA solutions. However, when looking at the evaluation of the respective contributions of physical and chemical mechanisms using our thermodynamic representation of the equilibrium,14 we note that the main difference between the systems {CO2 + MDEA + H2O} and {CO2 + TEA + H2O} comes from the difference in the energy released when the amine gets protonated. This seems to appear as the real cause of the diminution of the enthalpy of solution of CO2 in TEA solutions, compared to MDEA solutions. 5.3. Effect of the Hindrance on the Nitrogen Nucleus in the Case of Basic Ethanolamines for CO2 Absorption: Comparison among Mono-, Di-, and Triethanolamine. The comparison between CO2 dissolution in aqueous solution of MEA, DEA, and TEA will be based on experimental and calculated properties obtained at conditions of equilibrium between the unsaturated solution and CO2 gas phase at p = 0.5 MPa.

Figure 7. Enthalpy of solution (−(ΔsolH/(kJ·mol−1 of CO2)) versus CO2 loading for aqueous solution of TEA at T = 322.5 K and wTEA = 0.1500: ◇, experimental data at 1.0 MPa; thermodymanic model using: , (Ka) given by Simond et al.;22 −−, (Ka) given by Bates and Allen.21

that react with carbon dioxide to form mainly bicarbonates. Compared to MDEA, the substitution of a methyl group by an ethanol group leads to a decrease of pKa value for TEA. In these conditions, the limit of moles of CO2 absorbed per mol of amine (α) is lower in TEA solutions as observed in Figure 3. The difference in such solubility limits increases with temperature. On the other hand, this difference decreases when increasing pressure as the supplementary alcoholic group makes more hydrogen bonds possible and then facilitates physical 3594



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Figure 8. Enthalpy of solution (−(ΔsolH/(kJ·mol−1 of CO2)) versus CO2 loading for aqueous solution of TEA: ◇, experimental data at 1.0 MPa; from the thermodymanic model: (1) total enthalpy of solution; enthalpic contribution from reaction: (2) amine protonation (eq 10); (3) CO2 vapor−liquid equilibrium (eq 11); (4) second ionization of CO2 (eq 9); (5) first ionization of CO2 (eq 8). (a) T = 322.5 K and wTEA = 0.1500; (b) T = 322.5 K and wTEA = 0.3000; (c) T = 372.9 K and wTEA = 0.1500; (d) T = 372.9 K and wTEA = 0.3000.

Table 10. Enthalpies of Solution of CO2 (ΔsolH/kJ·molCO2−1) in Saturated Aqueous Solution of Mono-, Di-, and Triethanolamines (wTEA = 0.1500 and wTEA = 0.3000) at T = 322.5 K and T = 372.9 K for p = 0.5 MPa

Table 9. Limits of Solubility of CO2 Expressed as Loadings (α/molCO2·molamine−1) in Saturated Aqueous Solution of Mono-, Di- and, Triethanolamines (wTEA = 0.1500 and wTEA = 0.3000) at T = 322.5 K and T = 372.9 K for p = 0.5 MPa MEA4

DEA5 wTEAa

a

TEA

MEA4

0.956

0.595

0.647

MEA4

TEA

Tb = 322.5 K 0.963 0.652 Tb = 372.9 K 0.284 0.500

DEA5 wTEAa

wTEA = 0.3000

= 0.1500

0.769

DEA5 a

0.840

0.601

0.515

0.224

u(w) = 0.0001. bu(T) = 0.03 K.

a

The dissolution properties investigated are the limits of solubility, expressed as loading charge α (moles of CO2 dissolved per mole of amine) and, the enthalpies of solution (kJ·molCO2−1). All systems {CO2 + MEA + H2O},4 {CO2 + DEA + H2O},5 and {CO2 + TEA + H2O} were studied at T = 322.5 K and 372.9 K and compositions wamine = 0.1500 and 0.3000 (see Tables 9 and 10). The main difference when looking at the dissolution of carbon dioxide in aqueous solutions of TEA, DEA, and MEA results in the substitution on the nitrogen nucleus. The two latter amines, DEA (secondary amine) and MEA (primary amine), can react with carbon dioxide in aqueous solutions to form a carbamate. The stochiometry of the CO2 absorption would be thus restricted to a limiting loading α = 0.5 molCO2·molamine−1. However the carbamate formation can be

TEA

MEA4

74

90

73

TEA

wTEA = 0.3000

= 0.1500

90

DEA5 a

Tb = 322.5 K 47.1 56.9 Tb = 372.9 K 50.1 56.7

94

77

90

77

u(w) = 0.0001. bu(T) = 0.03 K.

reversed when increasing carbon dioxide pressure, and additional CO2 can be absorbed through acido-basic reaction with the amine restituted, leading to the formation of bicarbonate. According to speciation profiles derived from thermodynamic model representative of the liquid−vapor equilibria,5,23 molecular carbon dioxide in the unsaturated aqueous solutions of DEA and MEA could be neglected, as CO2 is mainly dissolved by chemical mechanisms with the formation of carbamates and bicarbonates. In the case of the dissolution of CO2 in aqueous solution of TEA, the total dissolved carbon dioxide corresponds to molecular CO2 and bicarbonate ions in solution. For all of these amine solutions, the conditions of pH at equilibrium are such that the quantity of CO2 absorbed with formation of carbonate could be neglected. 3595



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aqueous solution, a slight increase of the exothermic effect was observed. Despite the lack of solubility data for the system {CO2 + TEA + H2O}, which are necessary to regress interactions parameters for the thermodynamic model, we were able to estimate the enthalpy of solution of carbon dioxide into the absorbent by using interactions parameters from an amine from the same family (methyldiethanolamine). The thermodynamic model was then used to predict the enthalpy of solution in the same conditions as our experiments, testing two equilibrium constants for the amine protonation. Both thermodynamic constants were found to lead to the same results. At both temperatures (T = 322.5 K and T = 372.9 K) the amine protonation was found to provide the most important contribution with ∼ 80 % of the total energetic effect. For wTEA = 0.1500, the model gave good predictions for the enthalpy of solution. For wTEA = 0.3000, the predicted heat effects were found to be underestimated. The impact of the hindrance on the nitrogen nucleus as well as its degree of substitution on the enthalpy of solution were also discussed, comparing the dissolution of CO2 in aqueous solutions of MEA, DEA, TEA, and MDEA. Higher values of enthalpy of solution of CO2 in aqueous solutions of MDEA compared to TEA were attributed in the difference of energy released during the amine protonation. The differences in the energetic effect driving the absorption of carbon dioxide into aqueous solutions of mono-, di-, and triethanolamine were attributed to the formation of carbamate, which is favored for MEA compared to DEA and impossible for TEA. As carbamate formation is more energetic than bicarbonate formation, this results in higher values for the enthalpy of solution. In aqueous solutions of tertiary amine, part of the CO2 is dissolved following physical mechanism, less energetic than chemical absorption, resulting in a significantly reduced enthalpy of solution. As a final remark, this work showed that a thermodynamic model adjusted using only VLE properties could be used to represent the dissolution of CO2 in aqueous solutions of TEA over a wide range of temperature. However it was pointed out that the enthalpy of solution derived from such models is largely dependent on the accuracy of the speciation at equilibrium. Therefore direct measurements of enthalpies of solution are still of interest in order to test both the consistency of data as pointed out by Mathias and O'Connell24 and the ability of models to predict solution properties. Other data like experimental spectroscopic characterization of the solution for different loadings and/or computational simulations,25 would be very helpful to complete the other experimental data used to model the CO2 capture process and to get a better understanding of the role of structural differences between the different absorbents.

At our experimental conditions of pressure (p = 0.5 MPa) and temperature (T = 322.5 K, 372.9 K) the number of moles of carbon dioxide dissolved per mole of amine in aqueous solutions of DEA and MEA is superior to the stoichiometric limit (α = 0.5) that would be imposed if we were to consider only the formation of carbamates. In both case, we have to take into account carbamate hydrolysis, and the difference in solubility of CO2 between the two solutions can be thus explained by an easier inversion of the reaction of carbamate formation when using DEA. When comparing limits of solubility of CO2 in aqueous solutions of DEA and TEA, they appear to be similar at T = 322.5 K and wamine = 0.1500. The high solubility limit of CO2 (α = 0.956, T = 322.5 K, wTEA = 0.1500) observed for the latter is due to a combination of acido-basic reaction and physical dissolution. For this amine, physical dissolution is enhanced by the presence of three alcohol groups. As shown in Figure 9, the diminution of the pKa toward a value of ∼6.5 at T = 372.9 K leads to a dramatic decrease of the loading charge.

Figure 9. Temperature dependence of the pKa of MEA, DEA, and TEA22 from T = 323 K up to T = 373 K. , (C2H4OH)NH2; −−, (C2H4OH)2NH; −·−, (C2H4OH)3N.

Finally, regarding the impact on the energetic effect driving the absorption of carbon dioxide into aqueous solutions of mono-, di-, and triethanolamine, the highest enthalpy of solution (ΔsolH/kJ·molCO2−1) is observed in the aqueous solutions of MEA. Compared to DEA, MEA favors the dissolution of carbon dioxide under carbamate formation. As carbamate formation is more energetic than bicarbonate formation, this results in higher values for the enthalpy of solution. In TEA solutions, part of the CO2 is dissolved following physical mechanism, which is less energetic than chemical absorption, leading to a significantly reduced enthalpy of solution.

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6. CONCLUSION In this paper we have reported the first set of experimental data for enthalpy of solution for the system {CO2 + TEA + H2O}. Our solubility data determined from our enthalpy curves appeared to be consistent with literature studies using direct phase equilibria determination. At low loadings, the enthalpy of solution ΔsolH (kJ·molCO2−1) of carbon dioxide in aqueous solutions of TEA was observed to slightly decrease with pressure, and the diminution was attributed to change of residual enthalpies. A small increase of the energetic effect was also observed and could be attributed to the change of basicity of the amine with temperature. Regarding the composition dependency of the TEA

AUTHOR INFORMATION

Corresponding Author

*H.A.: Tel.: (00)1-519-824-4120 ext. 53811; fax: (00)1-519766-1499. E-mail: [email protected]. J.-Y.C.: Tel.: +33 (0)4 73 40 71 90; fax: +33 (0)4 73 40 53 28. E-mail: j-yves.coxam@ univ-bpclermont.fr. Funding

Financial support of this investigation by the environmental program PREVOIR, supported from Region Auvergne, is gratefully acknowledged. Notes

The authors declare no competing financial interest. 3596



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(20) Oscarson, J. L.; Wu, G.; Faux, W.; Izatt, R. M.; Christensen, J. J. Thermodynamics of protonation of alkanolamines in aqueous solution to 325 °C. Thermochim. Acta 1989, 154, 119−127. (21) Bates, R. G.; Allen, G. F. Acid Dissociation Constant and Related Thermodynamic Quantities for Triethanolammonium Ion in Water from 0 to 50 °C. J. Res. Natl. Bur. Stand. 1960, 64A, 343. (22) Simond, M. R.; Ballerat-Busserolles, K.; Coulier, Y.; Rodier, L.; Coxam, J.-Y. Dissociation Constants of Protonated Amines in Water at Temperatures from 293.15 to 343.15 K. J. Solution Chem. 2012, 41, 130−142. (23) Arcis, H. Etude Thermodynamique de la Dissolution du Dioxyde de Carbone dans des Solutions Aqueuses d′Alcanolamines. Ph.D. Thesis, Université Blaise Pascal, Clermont-Ferrand, France, 2009. (24) Mathias, P. M.; O’Connell, J. P. The Gibbs-Helmholtz Equation and the Thermodynamic Consistency of Chemical Absorption Data. Ind. Eng. Chem. Res. 2012, 51, 5090−5097. (25) Simond, M. R.; Ballerat-Busserolles, K.; Coxam, J.-Y., Padua, A. A. H. Molecular Simulations of Primary Alkanolamines Using an Extendable Force Field. ChemPhysChem 2012, DOI: 10.1002/ cphc.201200508.

ACKNOWLEDGMENTS The authors thank Sandrine Etien for her help in the literature search, Dr. Yohann Coulier for fruitful discussions on the modelization part, and Mickael Simond for providing images of the amines.

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Enthalpy of Solution of Carbon Dioxide in Aqueous Solutions of

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