Enthalpy of Solution of CO2 in Aqueous Solutions of Primary

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Enthalpy of Solution of CO2 in Aqueous Solutions of Primary Alkanolamines: A Comparative Study of Hindered and Nonhindered Amine-Based Solvents Hugues Arcis,*,† Yohann Coulier,‡,§ Karine Ballerat-Busserolles,‡,§ Laurence Rodier,‡,§ and Jean-Yves Coxam*,‡,§ †

Department of Chemistry, University of Guelph, Guelph, Ontario, Canada N1G 2W1 Clermont Université, Université Blaise Pascal, Institut de Chimie de Clermont-Ferrand, BP 10448, F-63000 Clermont-Ferrand, France § CNRS, UMR 6296, ICCF, BP 80026, F-63171 Aubière, France ‡

S Supporting Information *

ABSTRACT: Aqueous solutions of amines can be selective absorbents for the separation of carbon dioxide (CO2) from industrial effluents. Because of different chemical reactions, the mechanisms of CO2 absorption in 2-amino-2-methyl-1-propanol (hindered primary amine) or 2-aminoethanol (primary amine) will differ. These mechanisms are analyzed via the extents of reactions and enthalpies of solution, derived from a thermodynamic model representative of {CO2/amine/water} systems. The model was optimized to fit literature solubility data and tested by comparison of calculated and experimental enthalpy values. New experimental enthalpies of solution were collected for aqueous solutions of 2-amino-2-methyl-1-propanol (wamine = 0.1500, and wamine = 0.3000) at 372.9 K and pressures from 0.5 to 3.2 MPa.

1. INTRODUCTION Carbon capture is a possible approach to reducing the extent of anthropogenic release of greenhouse gas into the atmosphere and tempering climate change. CO2 absorption by amine solutions has been identified as a promising means of developing a CO2 capture process. Amine-based chemical solvents such as aqueous solutions of 2-aminoethanol (MEA) have been studied extensively for the development and optimization of the recovery of natural gas over the past 30 years. However, experimental conditions applied to CO2 capture require technologies with operating conditions different from those for the recovery of natural gas (lower temperature and pressure). Experimental thermodynamic or transport properties are lacking under the conditions of interest. Amine-based chemical solvents are known for their selective affinity in reacting with carbon dioxide, yielding a selective separation of CO2 from industrial exhaust fumes.1 Current industrial processes are based on absorption−desorption cycles in which the mechanism of gas absorption in such solutions combines physical dissolution and subsequent chemical reactions. Chemical reactions between carbon dioxide and the chemical solvent must be reversible to permit the regeneration of the absorbent. To reduce the energy cost of the desorption step, research of the amine selection has been conducted. Primary and secondary amines can react with CO2 to form carbamates,2 while classical acid−base reactions are observed between CO2 and tertiary amines. Depending on their classes, amines show different absorption properties, in general better kinetics for primary amines and lower CO2 desorption energies for tertiary amines.3 Hindered primary amines present intermediate absorption properties because of carbamate instability. © 2014 American Chemical Society

Thermodynamic models provide the theoretical tool needed to describe and understand the behavior of {CO2/amine/H2O} systems, particularly for identifying the chemical mechanisms responsible for CO2 absorption. The development of such models requires knowledge of the thermodynamic properties characteristic of equilibria such as constants of reaction (K), Henry’s constants (H), and vapor (p) and partial (pCO2, pH2O, and pamine) pressures. Once developed, these models offer powerful insights into the system by providing an estimation of phase speciation at equilibrium4,5 through the extent (ξ) of the reactions involved in the CO2 absorption mechanism. Such models that are representative of vapor−liquid equilibria are usually optimized to fit solubility data.6−9 The prediction of chemical speciation will then rely on several parameters (accuracy of the solubility data, chemical reactions presumed, choice of activity, and fugacity coefficient models). Derivative properties such as the enthalpy of solution (ΔsolH) depend strongly on the estimated speciation, and direct comparison with experimental values constitutes a severe test of model consistency.4,5 The objective of this work was to compare the chemical mechanisms involved in the absorption of CO2 when using different primary amine-based solvents, 2-amino-2-methyl-1propanol (AMP), a hindered amine, and 2-aminoethanol (MEA), a nonhindered amine. The comparison relies on energy effect differences observed between the two systems. Received: Revised: Accepted: Published: 10876

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Figure 1. Schematic of the experimental setup for measuring heats of absorption.

CO2 molar flow rate to the total amine molar flow rate. The heat power of absorption divided by molar flow rates of CO2 or amine gives the enthalpy of solution per mole of CO2 (ΔsolH, in kilojoules per mole of CO2) or per mole of amine (ΔsolH, in kilojoules per mole of AMP). The total molar flow rates are obtained from the pump flow rates, using densities, mass compositions of solution, and molar masses of CO2, amine, and water. The densities of aqueous AMP solutions were taken from ref 11. The densities of CO2 were calculated with ALLPROPS.12 The uncertainty of the enthalpy data is estimated to be between 1 and 3% and to reach 5% for the smallest measured heat of mixing. 2.3. Chemicals and Materials. The 2-amino-2-methyl-1propanol was obtained from Sigma-Aldrich (>97% pure) and used without further purification. Carbon dioxide (99.998% pure) was obtained from Saga. Water was distilled and degassed before being used (resistivity of 18.2 MΩ cm). Aqueous solutions were prepared by mass; uncertainty in the mass fraction (w) was estimated to be less than ±10−4. Aqueous solutions were stored in glass bottles in an opaque cabinet to prevent any photodegradation.

The enthalpies of solution (ΔsolH) of CO2 in both absorbents were investigated at 322.5 and 372.9 K and at pressures from 0.5 to 5 MPa. Some of the experimental data were taken from previous studies,10,11 while additional measurements for the {CO2/AMP/H2O} system at 372.9 K and pressures from 0.5 to 3 MPa are reported here. Experimental data were used to validate a γ−ϕ thermodynamic model adjusted for vapor liquid equilibrium data. Results derived from the model, such as gas solubilities, extents of reaction, and enthalpies of solution, were used to discuss the different mechanisms involved in carbon dioxide absorption in aqueous solutions of AMP and MEA.

2. EXPERIMENTAL SECTION 2.1. Experimental Arrangement. The experimental flow calorimetric technique has been described previously.10 The enthalpy of solution of CO2 in the absorbent was measured using a custom-made flow-mixing cell adapted to a Setaram C-80 heat conduction differential calorimeter. A schematic of the experimental setup is depicted in Figure 1. Experiments were conducted at a constant temperature (±0.01 K) and pressure (±0.02 MPa) as a function of the gas loadings, α (moles of CO2 per mole of amine). Two high-pressure syringe pumps (Isco model 100 DM) were used to inject both CO2 and the absorbent into the mixing cell. The temperature of the injected fluids before they enter the calorimeter was adjusted carefully and precisely using a system of preheaters.10 The system was maintained at a constant pressure with a pressure regulator located at the end of the flow line. The pressure was measured along the flow line before and after the mixing cell with the help of three pressure gauges. 2.2. Operating Procedure and Experimental Uncertainty. The heat power of absorption is determined at different CO2 loadings (α). Loadings (α) are set by the ratio of the total

3. EXPERIMENTAL RESULTS Enthalpies of solution of CO2 in aqueous solutions of AMP were measured at 372.9 K and 0.5, 1, or 3 MPa, and for two absorbent compositions (wAMP = 0.1500 and 0.3000). Experiments were conducted at different loadings (α), up to the saturation of the absorbent solution. The experimental data are presented in the Supporting Information. Enthalpy data were used to derive the CO2 limit of solubility into the absorbents. As shown in Figure 2 when the solution is unsaturated, the enthalpy of solution, ΔsolH, expressed in joules per mole of AMP increases with α until it reaches a maximal 10877

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Figure 2. Enthalpies of solution of CO2 in aqueous AMP solutions (−ΔsolH, in kilojoules per mole of AMP) vs CO2 loadings at 372.9 K and (□) 0.5, (◇) 1.0, and (○) 3.2 MPa. (a) w = 0.1500; (b) w = 0.3000.

Table 1. Experimental Values of CO2 Loadings at Saturation in Aqueous AMP Solutions (w = 0.1500 and 0.3000) at 372.9 K w = 0.1500

w = 0.3000

p (MPa)

αsat (mol of CO2/ mol of amine)

δαsat

p (MPa)

αsat (mol of CO2/ mol of amine)

δαsat

0.55 1.09 3.20

0.814 0.962 1.11

0.02 0.03 0.03

0.58 1.06 3.14

0.688 0.858 1.00

0.02 0.03 0.03

value that remains more or less constant. This plateau, characteristic of a saturated solution, can in some case slightly decrease because of vaporization effects in the CO2 gas phase, which has already been observed experimentally for others systems.10 The intersection between unsaturated (enthalpy increase) and saturated (plateau) domains yields the gas loading (αsat) at saturation or the solubility limit (s). Experimental gas loadings at saturation (αsat) in aqueous AMP solutions (wAMP = 0.1500 and 0.3000) were graphically determined at 372.9 K; values are listed in Table 1 together with experimental uncertainties.

4. VAPOR LIQUID EQUILIBRIUM (VLE) THERMODYNAMIC MODEL The thermodynamic model was derived to describe vapor liquid equilibria (VLE) using a γ−ϕ approach. The model is similar to that previously used to represent the absorption of CO2 in aqueous solutions of N-methyldiethanolamine (MDEA)4 and diethanolamine (DEA).13 Briefly, the modeling is based on a system of equations representative of chemical equilibra (eqs 1−5) and vapor−liquid equilibria involved in the absorption process. H 2O ⇌ H+ + OH−

(1)

CO2 (aq) + H 2O ⇌ HCO3− + H+

(2)

HCO3− ⇌ CO32 − + H+

(3)

RNH 2 + H+ ⇌ RNH3+

(4)

RNH 2 + HCO3− ⇌ RNHCOO− + H 2O

(5)

Figure 3. Experimental conditions (○) (T, p, and mamine) for VLE literature data used to adjust Pitzer interaction parameters (eq 7).

log K = A +

B + C × ln T + DT T

(6)

where A−D are empirical parameters listed in the Supporting Information. With regard to MEA and AMP protonation (eq 4), parameters used in eq 6 were derived by fitting literature equilibrium constants.16−28 Because of experimental challenges associated with their determination, experimental equilibrium constants for carbamate formation (eq 5) are scarce. The only temperature-dependent data currently available in the literature have been derived via an NMR16,23 technique, determined as adjusted parameters in thermodynamic models,9,24,25 or calculated using computational chemistry,26,28 yielding poor accuracy in the determination of the equilibrium constants. For MEA, McCann et al.27 measured the carbamate equilibrium constant with a calorimetric technique but only at 298 K. We chose to use the constants reported by Bottinger et al.16 because they were

Chemical equilibra (eqs 1−5) are defined by their equilibrium constant, K. Equilibrium constant Kw for water dissociation (eq 1) was taken from ref 14. The first and second dissociation constants of CO2 (eqs 2 and 3, respectively) were obtained from ref 15. For both amines (AMP and MEA), equilibrium constants of amine protonation (eq 4) and carbamate formation (eq 5) were calculated using eq 6: 10878

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Figure 4. CO2 loading at saturation (αsat) in aqueous MEA and AMP solutions vs total pressure. (a) wMEA = 0.1500: (□) 322.5 K10 and (○) 372.9 K.10 (b) wAMP = 0.1500: (△) 322.5 K11 and (◇) 372.9 K (this work). (c) wMEA = 0.3000: (■) 322.5 K10 and (●) 372.9 K.10 (d) wAMP = 0.3000: (▲) 322.5 K11 and (◆) 372.9 K (this work). The solid lines are predictions from the thermodynamic model.

Figure 5. Enthalpies of solution (ΔsolH, in kilojoules per mole of CO2) vs CO2 loading for aqueous solutions of AMP. (◇) Experimental data (for 322.5 K and 0.5 MPa11 or for 372.9 K and 1 MPa). From the thermodynamic model: (1) total enthalpy of solution. From the enthalpic contribution from reaction: (2) amine protonation (eq 4), (3) carbamate formation (eq 5), (4) CO2 vapor−liquid equilibrium, (5) first ionization of CO2 (eq 2), and (6) second ionization of CO2 (eq 3). (a) T = 322.5 K, and w = 0.1500. (b) T = 322.5 K, and w = 0.3000. (c) T = 372.9 K, and w = 0.1500. (d) T = 372.9 K, and w = 0.3000.

Figure 6. Enthalpies of solution (ΔsolH, in kilojoules per mole of CO2) vs CO2 loading for aqueous solutions of MEA. (◇) Experimental data at 1 MPa.10 From the thermodynamic model: (1) total enthalpy of solution. From the enthalpic contribution from reaction: (2) amine protonation (eq 4), (3) carbamate formation (eq 5), (4) CO2 vapor−liquid equilibrium, (5) first ionization of CO2 (eq 2), and (6) second ionization of CO2 (eq 3). (a) T = 322.5 K, and w = 0.1500. (b) T = 322.5 K, and w = 0.3000. (c) T = 372.9 K, and w = 0.1500. (d) T = 372.9 K, and w = 0.3000.

derived from the experimental study covering the widest temperature range; the data from the others sources summarized in

the Supporting Information were found to be consistent with the results of Bottinger et al.16 In the case of AMP, the carbamate 10879

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Figure 7. Speciation at equilibrium for the {CO2/MEA/H2O} system derived from the thermodynamic model: (···) MEA, () MEAH+, (−··−) MEACOO−, (−−−) HCO3−, (---) CO32−, and (−·−) CO2(aq). (a) T = 322.5 K, and wMEA = 0.1500. (b) T = 322.5 K, and wMEA = 0.3000. (c) T = 372.9 K, and wMEA = 0.1500. (d) T = 372.9 K, and wMEA = 0.3000.

⎤ ⎡ I 2 ln γi = −Aϕzi 2⎢ + ln(1 + b I )⎥ ⎦ ⎣1 + b I b ⎧ ⎫ βi1, j ⎪ ⎪ + 2 ∑ mj⎨βi0, j + [1 − (1 + 2 I ) exp( −2 I )]⎬ ⎪ ⎪ 2I j≠w ⎩ ⎭

is known to be unstable, yielding very small equilibrium constants of formation over the temperature range of interest. Sartori et al.29 analyzed aqueous solutions of AMP charged with CO2 by 13C NMR and concluded that AMP reacts with CO2 to mainly form protonated amine (AMPH+) and bicarbonate (HCO3−). However, Cifta et al.30 found experimental evidence of carbamate formation at low loadings. Recently, Fernandes et al.23 conducted a systematic study of carbamate formation for several primary and secondary amines using 1H NMR spectroscopy but reported that AMP carbamates were found to be too unstable to be detected. Direct experimental determination of accurate equilibrium constants for the formation of AMP carbamate is difficult, as suggested by the large scatter observed among the rare experimental studies available. Another approach consists of considering the equilibrium constant as an adjustable parameter in the VLE thermodynamic model. Using a γ−ϕ model, Xu et al.9 derived equilibrium constants decreasing from 0.03 at 313 K to 8.9 × 10−3 at 373 K. Using a simplified model,31 Park et al.25 obtained values decreasing from 0.27 at 313 K to 3.0 × 10−8 at 353 K. More recently, Mehdizadeh et al.28 reported constants using theoretical computational chemistry that decrease from 4.6 at 313 K to 1.75 at 373 K. Coefficients in eq 6 for the constant of carbamate formation of AMP were adjusted to the literature data reported in the Supporting Information. A Pitzer model32 and a virial equation of state were used to describe the nonideality in the liquid and gas phases, respectively. More details can be found elsewhere.4,13 The activity coefficients, γ, in the liquid phase were calculated with eq 7:

− zi

2

∑ ∑ mjmk j≠w k≠w

exp( − 2 I )]

βj1, k 4I 2

[1 − (1 + 2 I + 2I ) (7)

where mi and zi are the molality and charge of species i, respectively; mj and mk are the molalities of species j and k, respectively; I is the ionic strength in moles per kilogram; Aϕ, in units of (kg/mol)−1/2, is the Debye−Hückel limiting slope for the osmotic coefficient at the experimental temperature and pressure; β0i,j and β1i,j are binary Pitzer interaction parameters; and Pitzer parameter b = 1.2 kg1/2 mol−1/2. We followed the approach adopted by Oscarson et al.33 and considered only binary interaction parameters, β0i,j and β1i,j, classified into four groups: neutral−neutral, ion−neutral, cation− anion, and like charged ions. Interactions between ions of the same sign were all assigned to zero. Interaction parameters were obtained through an optimization process to fit solubility data;8,24,34−50 values are reported in the Supporting Information. The model was adjusted by fitting the total CO2 absorbed, expressed as the partial pressure or total pressure, for different temperatures and absorbent compositions. We selected 379 and 348 experimental VLE data points for the {CO2/AMP/H2O} and {CO2/MEA/H2O} systems, respectively, to optimize 10880

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Figure 8. Speciation at equilibrium for the {CO2/AMP/H2O} system derived from the thermodynamic model: (···) AMP, () AMPH+, (−··−) AMPCOO−, (−−−) HCO3−, (---) CO32−, and (−·−) CO2(aq). (a) T = 322.5 K, and wAMP = 0.1500. (b) T = 322.5 K, and wAMP = 0.3000. (c) T = 372.9 K, and wAMP = 0.1500. (d) T = 372.9 K, and wAMP = 0.3000.

parameters β0i,j and β1i,j with standard deviations of 0.36 and 0.24 MPa, respectively. Experimental conditions of temperature, pressure, and absorbent composition expressed as amine molality are given in the Supporting Information and shown in Figure 3.

to solvate CO2 when the amine composition (w) is increased, while the second effect (temperature) is related to the change in amine basicity with temperature. At the same temperature and amine composition, the absorption (αsat) is much higher in aqueous solutions of AMP than in those of MEA. The absorption difference between both solvents increases with temperature, which is consistent with temperature dependence differences for both amine pKa values.19,22 The higher pKa value for AMP is explained by the substitution of the ethanol group with a methyl-propanol group. Also, because of its steric hindrance, the AMP carbamate is unstable. For this reason, CO2 absorption mechanisms are more likely to be similar to those observed with a tertiary amine-based solvent, yielding higher loadings at saturation compared to MEA solutions. Differences in absorption mechanisms between both solvents are discussed in more detail below with the help of speciation calculation from the thermodynamic model. 5.2. Enthalpy Data. The thermodynamic model was adjusted to fit the CO2 partial pressure as a function of temperature and absorbent composition. Enthalpies of solution were derived from the model4,13 and compared with experimental values to test its consistency. Comparisons between modeling and experimental data are shown Figures 5 and 6 for AMP and MEA, respectively. A typical experimental trend for nonhindered primary and secondary amines is that the enthalpy of solution remains constant at low loadings (α), up to α ≈ 0.5. Then, ΔsolH decreases because of carbamate hydrolysis. For the tertiary amine, ΔsolH remains constant up to the solubility limit (αsat) of CO2 in the solution. Our thermodynamic model reproduces qualitatively these experimental trends for MEA solutions (nonhindered amine),

5. RESULTS AND DISCUSSION 5.1. Solubility Data. The model described in section 4 was used to predict solubility data at different temperatures and absorbent compositions as a function of loading. Temperatures and amine concentrations were chosen to match our experimental conditions. Results for the {CO2/AMP/H2O} and {CO2/MEA/H2O} systems are shown in Figure 4. Below 3 MPa, we observed a reasonable agreement between experimental and calculated αsat values for both AMP and MEA solutions. At the highest pressures, the experimental solubility values are lower than those predicted by the model because of the difficulties in estimating graphically the α value corresponding to the first enthalpy of solution on the plateau (saturated domain) in Figure 2. When carbon dioxide is chemically absorbed, the enthalpy of solution increases linearly and the limit of solubility corresponds to the intersection with the plateau. As loading α increases, carbon dioxide becomes partially absorbed because of physical mechanisms. Physical absorption is less energetic than chemical absorption, leading the enthalpy of solution also to increase but not linearly. Solubilities of CO2 in the aqueous solutions of AMP and MEA follow the same trends: a decrease with amine composition and temperature. The first effect (amine composition) can be explained by the diminution of water molecules available 10881

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Figure 9. Fraction of CO2 absorbed in solution at equilibrium for the {CO2/MEA/H2O} system derived from the thermodynamic model: (−··−) MEACOO−, (−−−) HCO3−, (---) CO32−, and (−·−) CO2(aq). (a) T = 322.5 K, and wMEA = 0.1500. (b) T = 322.5 K, and wMEA = 0.3000. (c) T = 372.9 K, and wMEA = 0.1500. (d) T = 372.9 K, and wMEA = 0.3000.

studies of excess properties, following the approaches of Orozco et al.51 or Simonds et al.,52 and the acquisition of experimental speciation data. 5.3. Differences in the Absorption Mechanisms between Hindered and Nonhindered Amine-Based Solvents. Chemical reactions involved in the overall absorption of carbon dioxide, ΔsolH, yield different enthalpies of reaction (eqs 1−5). Details of the different contributions are shown in Figures 5 and 6 for the {CO2/AMP/H2O} and {CO2/MEA/ H2O} systems. At the same temperatures, enthalpies of solution (ΔsolH) per mole of CO2 are higher in solutions of MEA than in solutions of AMP. It has been shown that the amine concentration does not affect substantially the enthalpy of solution of CO2.11 The strongest effect arises from amine protonation (eq 4), where it is more exothermic when AMP is the absorbent (Figure 5) than when MEA is (Figure 6). The greater enthalpy of solution in aqueous solutions of MEA compared to AMP is explained by the difference observed in the enthalpy of carbamate formation, contributing to ∼30% of the total energy of absorption for MEA and being almost negligible for AMP. It is now interesting to consider the chemical mechanisms involved in the absorption of CO2 to understand differences between hindered and nonhindered amine-based solvents. Carbon dioxide reacts with two primary amine molecules to form a carbamate according to the mechanism (eq 8) proposed by Caplow,53

but also when AMP is the solvent. AMP is a hindered primary amine that shows intermediate behavior between that of a nonhindered primary amine and that of a tertiary amine. The reliability of the model is shown in Figures 5 and 6, where good agreement is observed between experimental and calculated enthalpies of solution. The agreement is excellent for the MEA solutions at all temperatures within ±5% deviation. For the AMP solutions, good agreement is observed at 373 K. At 323 K, the model overestimated the energetic effect, especially at low loadings where the deviation is ∼15%. A recent study by Mehdizadeh et al.28 reported the enthalpy of absorption of CO2 in 25 wt % AMP at 313, 353, and 393 K. A direct comparison with our experimental data proved to be difficult as no numerical values were reported in their paper. However, their Figure 628 shows that their experimental enthalpy data, scattering between 80 and 95 kJ/mol of CO2 at low loadings, are consistent with our experimental values. Difficulties in modeling CO2 absorption in aqueous solutions of AMP, especially at low loadings, has been already mentioned by Mehdizadeh et al.28 and Dash et al.40 The thermodynamic model shows that CO2 solubilizes in AMP solutions mostly to form bicarbonate (HCO3−). While the equilibrium constant of the first dissociation of CO2 (eq 2) is known accurately, small inconsistencies for equilibrium constants of amine protonation (eq 4) and carbamate formation (eq 5) can affect the model. The modeling may also be influenced by the choice of activity model and inconsistency in the solubility database used to adjust its semiempirical interaction parameters. Estimation of activity coefficients could be improved by the development of theoretical

CO2 + 2RNH 2 ⇌ RNHCOO− + RNH3+ 10882

(8)

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Figure 10. Fraction of CO2 absorbed in solution at equilibrium for the {CO2/AMP/H2O} system derived from the thermodynamic model: (−··−) AMPCOO−, (−−−) HCO3−, (---) CO32−, and (−·−) CO2(aq). (a) T = 322.5 K, and wAMP = 0.1500. (b) T = 322.5 K, and wAMP = 0.3000. (c) T = 372.9 K, and wAMP = 0.1500. (d) T = 372.9 K, and wAMP = 0.3000.

at lower loadings (α < 0.5) and a low temperature (322.5 K) because of pH conditions and the high pKa value. The amount of carbonate goes through a maximum when α = 0.5, according to the calculated speciation in Figure 8. At higher temperatures, it remains small compared to that of bicarbonate formation but appears to be favored as the amine concentration in the solvent is increased (Figure 10). In conclusion, the greatest exothermic effect observed in solutions of MEA compared to those of AMP is due to the ability of MEA to form a stable carbamate, yielding a significant additional exothermic effect, while for AMP this effect is almost negligible.

limiting CO2 absorption to a maximal loading (αsat) of 0.5 (loading at saturation). However, αsat can exceed this value because of possible carbamate hydrolysis following eq 9: RNHCOO− + H 2O ⇌ HCO3− + RNH 2

(9)

Hindered amines are known to generate unstable carbamates, and when such amine-based chemical solvents are used, carbon dioxide will be absorbed in solution as bicarbonates and carbonates. Our model provides a means of representing both qualitatively and quantitatively the {CO2/amine/H2O} system and discussing mechanisms involved in CO2 absorption. CO2 absorption involves several chemical reactions, and at thermodynamic equilibrium, the concentration of each species present in solution depends on gas loading α and temperature. Figures 7 and 8 show speciation curves derived from our model, and Figures 9 and 10 show what form CO2 is absorbed in solution as a function of the loadings. Our results support the mechanism discussed above (eqs 8 and 9). As shown in Figures 7 and 9, CO2 is absorbed in MEA solutions (nonhindered primary amine) with formation of bicarbonate (HCO3−) and carbamates (MEACOO−) at loadings (α) of ≤0.5. As the loading is increased over 0.5, carbamates start to progressively be hydrolyzed, restoring the amine that will react with CO2 and form bicarbonates. The formation of carbonates appears to be negligible under all experimental conditions. In the case of AMP (hindered primary amine), speciation curves (Figures 8 and 10) show no evidence of significant carbamate formation. Carbon dioxide reacts with AMP to form bicarbonates and carbonates. The formation of carbonates is found to be significant

6. CONCLUSION CO2 absorption in primary amine-based solvents was studied by solution calorimetry at 322.5 and 372.9 K. We reported new enthalpies of solution for the {CO2/AMP/H2O} system measured at 372.9 K, pressures of 0.5, 1, and 3 MPa, and wAMP values of 0.1500 and 0.3000. A rigorous thermodynamic model was proposed to describe the solution chemistry for CO2 absorption in a primary amine-based solvent with both regular (MEA) and hindered (AMP) amines. The model is representative of chemical and physical equilibria involved in the absorption process and was used to derive enthalpies of solution. The reliability of the model was tested by comparing calculated enthalpies to experimental data. A good agreement was found for CO2 absorption in aqueous MEA solutions with a mechanism based on carbamate formation. The model can predict the enthalpy of solution to ±5%, showing its aptitude in providing an accurate estimation of the speciation in solution. In AMP solutions, the model showed a similar predictive ability 10883

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at 372.9 K, but at 322.5 K, deviations of up to ±15% with experimental enthalpies were observed. However, this model provided enough foundation to discuss qualitatively the mechanisms involved in CO2 absorption in an AMP solution. It was found that because of the instability of carbamate formed with AMP, CO2 is absorbed in solution as bicarbonate and carbonate. Because of its high pKa value, significant formation of carbonates was observed. It has been demonstrated that rigorous thermodynamic models can be a strong theoretical asset for describing qualitatively and quantitatively the mechanisms involved in CO2 absorption. Accurate experimental speciation data from conductivity, potentiometric, or spectrometric techniques would help to improve these models, especially data at low loadings (α < 0.5). Finally, this work provides some of the information required for developing engineering tools needed to characterize the optimal operating conditions in CO2 capture processes. However, such tools need to include others parameters such as kinetics of reaction, solvent degradation, corrosion factors, etc.

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(7) Aroua, M. K.; Haji-Sulaiman, M. Z.; Ramasamy, K. Modelling of Carbon Dioxide Absorption in Aqueous Solutions of AMP and MDEA and Their Blends Using Aspenplus. Sep. Pur. Technol. 2002, 29, 153− 162. (8) Kundu, M.; Mandal, B. P.; Bandyopadhyay, S. S. Vapor Liquid Equilibrium of CO2 in Aqueous Solutions of 2-Amino-2-methyl-1propanol. J. Chem. Eng. Data 2003, 48, 789−796. (9) Xu, S.; Wang, Y. W.; Otto, F. D.; Mather, A. E. Representation of the Equilibrium Solubility Properties of CO2 with Aqueous Solutions of 2-Amino-2-Methyl-1-Propanol. Chemical Engineering and Processing: Process Intensification 1992, 31, 7−12. (10) Arcis, H.; Ballerat-Busserolles, K.; Rodier, L.; Coxam, J. Y. Enthalpy of Solution of Carbon Dioxide in Aqueous Solutions of Monoethanolamine at Temperatures of 322.5 and 372.9 K and Pressures up to 5 MPa. J. Chem. Eng. Data 2011, 56, 3351−3362. (11) Arcis, H.; Rodier, L.; Coxam, J. Y. Enthalpy of Solution of CO2 in Aqueous Solutions of 2-Amino-2-methyl-1-propanol. J. Chem. Thermodyn. 2007, 39, 878−887. (12) ALLPROPS. Thermodynamic Properties of Fluids, version 4.1; Center for Applied Thermodynamics, University of Idaho: Moscow, ID, 1995. (13) Arcis, H.; Ballerat-Busserolles, K.; Rodier, L.; Coxam, J.-Y. Measurement and Modeling of Enthalpy of Solution of Carbon Dioxide in Aqueous Solutions of Diethanolamine at Temperatures of (322.5 and 372.9) K and Pressures up to 3 MPa. J. Chem. Eng. Data 2012, 57, 840−855. (14) Marshall, W. L.; Franck, E. U. Ion Product of Water Substance, 0−1000 °C, 1−10,000 bar New International Formulation and Its Background. J. Phys. Chem. Ref. Data 1981, 10, 295−304. (15) Renon, H. Representation of NH3-H2S-H2O, NH3-SO2-H2O, and NH3-CO2-H2O Vapor-Liquid Equilibria. In Thermodynamics of Aqueous Systems with Industrial Applications; American Chemical Society: Washington, DC, 1980; Vol. 133, pp 173−186. (16) Bottinger, W.; Maiwald, M.; Hasse, H. Online NMR Spectroscopic Study of Species Distribution in MEA−H2O−CO2 and D EA−H2O−CO2. Fluid Phase Equilib. 2008, 263, 131−143. (17) Littel, R. J.; Bos, N.; Knoop, G. J. Dissociation Constants of Some Alkanolamines at 293, 303, 318, and 333 K. J. Chem. Eng. Data 1990, 35, 276−277. (18) Hamborg, E. S.; Versteeg, G. F. Dissociation Constants and Thermodynamic Properties of Amines and Alkanolamines from (293 to 353) K. J. Chem. Eng. Data 2009, 54, 1318−1328. (19) 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. (20) Bates, R. G.; Pinching, G. D. Acidic Dissociation Constant and Related Thermodynamic Quantities for Monoethanolammonium Ion and Water from 0 to 50 °C. J. Res. Natl. Bur. Stand. (U.S.) 1951, 46, 349−352. (21) Datta, S. P.; Grzybowski, A. K. 593. Acid Dissociation Constants of the Ammonium Group in 2-Aminoethanol, 2-Aminoethyl Phosphate, and 2-Aminoethyl Sulphate. J. Chem. Soc. 1962, 3068− 3077. (22) Ballerat-Busserolles, K.; Simond, M. R.; Coulier, Y.; Coxam, J. Y. Protonation of Alkanolamines and Cyclic Amines in Water at Temperatures from 293.15 to 373.15 K. Pure Appl. Chem. 2014, 86, 233−243. (23) Fernandes, D.; Conway, W.; Burns, R.; Lawrance, G.; Maeder, M.; Puxty, G. Investigations of Primary and Secondary Amine Carbamate Stability by 1H NMR Spectroscopy for Post Combustion Capture of Carbon Dioxide. J. Chem. Thermodyn. 2012, 54, 183−191. (24) Jou, F.-Y.; Mather, A. E.; Otto, F. D. The Solubility of CO2 in a 30 Mass Percent Monoethanolamine Solution. Can. J. Chem. Eng. 1995, 73, 140−147. (25) Park, S. H.; Lee, K. B.; Hyun, J. C.; Kim, S. H. Correlation and Prediction of the Solubility of Carbon Dioxide in Aqueous Alkanolamine and Mixed Alkanolamine Solutions. Ind. Eng. Chem. Res. 2002, 41, 1658−1665.

ASSOCIATED CONTENT

S Supporting Information *

Experimental enthalpies of solution of CO2 in aqueous solutions of AMP (T = 372.9 K; p = 0.5, 1, or 3 MPa; and wAMP = 0.1500 and 0.3000), parameters used in eq 6 to represent the temperature dependence of equilibrium constants for eqs 4 and 5; coefficients for binary Pitzer interaction parameters used in eq 7; and VLE data sources and experimental conditions used to adjust those parameters. This material is available free of charge via the Internet at http://pubs.acs.org.

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AUTHOR INFORMATION

Corresponding Authors

*E-mail: [email protected]. Telephone: (519) 824-4120, ext. 53811. Fax: (519) 766-1499. *E-mail: [email protected]. Telephone: +33 (0)4 73 40 71 90. Fax: +33 (0)4 73 40 53 28. Notes

The authors declare no competing financial interest.

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REFERENCES

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