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Heats of Absorption of CO2 in Aqueous Solutions of Tertiary Amines: N‑Methyldiethanolamine, 3‑Dimethylamino-1-propanol, and 1‑Dimethylamino-2-propanol Aravind V. Rayer and Amr Henni* Acid Gas Removal Laboratory, Faculty of Engineering and Applied Science, University of Regina, Regina, Saskatchewan, Canada S4S 0A2 S Supporting Information *

ABSTRACT: The heats of absorption of CO2 in aqueous tertiary amines (N-methyldiethanolamine (MDEA), 3-dimethylamino1-propanol (3DMAP), and 1-dimethylamino-2-propanol (1DMAP)) were measured using a flow calorimeter at different temperatures (298.15, 313.15, and 343.15 K) and pressures (2 and 5 MPa) for 5 and 30 wt %. The integral heat of absorption and the indirect solubility limit were determined using the experimental data. The heat of absorption values were correlated using a rigorous thermodynamic model. The temperature dependent heats of reactions and the equilibrium constants for amine protonation reactions were calculated using the model. Comparisons were made with the integral heat of absorption values obtained from phase equilibrium measurements for MDEA using the Gibbs−Helmholtz equation. 3DMAP was found to have a higher CO2 loading capacity (molCO2/molamine) and a higher reaction rate with a similar heat of absorption when compared to other popular tertiary amines (MDEA and triethanolamine (TEA)).

1. INTRODUCTION Absorption/stripping with aqueous amines is considered to be a technically viable option for capturing CO2 from natural gas, flue gas in petroleum refining, coal gasification, and hydrogen production. However, the cost of CO2 capture using aqueous solutions of a single amine is considered to be quite high: ≈40−70 US$/ton of CO2.1 The industrial process is based on a combination of chemical reaction and physical dissolution in aqueous amine solutions using an absorber and a regeneration phase that is done in a stripper unit. Eighty percent of the operational cost is the cost of regeneration of the solvent. There is, therefore, a need to develop solvents that require lower energy for regeneration. However, other factors such as solvent capacity, corrosion potential, and solvent stability along with better absorption and regeneration characteristics must also be taken into account. Tertiary amines have lower regeneration energy but also a lower absorption rate; meanwhile primary and secondary amines are effective in the absorption process, but are more difficult to regenerate because of the high heat of reactions involved. The cost of steam required in the amine regeneration stage represents half the operating cost of a plant, and depends mainly on the design of unit operations for acid gas removal. The heat of reaction between CO2 and amine and the heat of dissolution are important factors in determining the thermal effects observed in gas treating processes. The combination of these heats is called the heat of absorption. Heat of absorption values are measured directly by calorimetry or estimated from vapor−liquid equilibrium data using the Gibbs−Helmholtz thermodynamic relation. The uncertainty of this thermodynamic relation is about 20−30% for the most accurate solubility data.2 This estimation does not explain the alteration of the process due to the temperature difference, because the © 2014 American Chemical Society

estimated enthalpies are differential in loading but integral in temperature. On the other hand, experimentally measured enthalpies using calorimeters are integral in loading but differential in temperature, and they represent all the heat released from mixing the solution and acid gas from zero loading to the final loading. Therefore, direct calorimetric measurements are necessary to develop consistent thermodynamic models for the correlation and the prediction of both solubilities and enthalpies of solutions for the CO2 + H2O + amine systems. Batch calorimetry and flow calorimetry are two techniques used to measure the heat of absorption. The heat released by introducing an amount of acid gas in the solvent is measured by batch calorimetry,3 and the heat released, by feeding simultaneously acid gas and solvent at different flow rates for the desired acid gas loading, is measured by flow calorimetry.4 Primary and secondary amines have high reactivity with CO2 and tend to have high heats of reaction (≈80 kJ·molCO2−1) due to the formation of a carbamate complex. The extent of heat required to break down this complex during the regeneration stage depends mainly on the heat of reaction. Tertiary amines with lower heat of reaction (≈60 kJ·molCO2−1) require less energy for regeneration as CO2 is absorbed primarily as a bicarbonate. Tertiary amines, in many applications, must be promoted by primary or other fast-reacting amines to provide adequate rates of CO2 mass transfer. Tertiary amines seem to be less inclined to oxidative degradation and do not participate in degradation by carbamate polymerization.5 The order of Received: Revised: Accepted: Published: 4953

December 6, 2013 February 23, 2014 February 24, 2014 February 24, 2014 dx.doi.org/10.1021/ie4041324 | Ind. Eng. Chem. Res. 2014, 53, 4953−4965

Industrial & Engineering Chemistry Research

Article

Table 1. Summary of Previous Work on the Heats of Absorption of CO2 in Aqueous Tertiary Amine Systems

a

T/K

authors

apparatus

solvent

Merkley et al.9 Oscarson et al.37 Mathonat et al.4 Kierzkowska-Pawlak and Zarzycki38 Carson et al.29 Acris et al.7 Kim and Svendsen39

in-house isothermal flow calorimeter in-house isothermal flow calorimeter Setaram C-80 flow calorimeter CPA-102 differential calorimeter in-house displacement calorimeter in-house Dewar-type calorimeter CPA-122 differential calorimeter

Rodier et al.8

Setaram C-80 flow calorimeter

MDEA (20, 40, 60 wt %) MDEA (20, 35, 50 wt %) MDEA (30 wt %) MDEA (10, 20, 30 wt %) MDEA (10, 20, 30 wt %) MDEA (15, 30 wt %) MDEA (30, 40 wt %) DEEAa (32, 37 wt %) DMAEOEa (26.3 wt %) 3DMAP (26.3 wt %)

288.7, 299.9, 313.2, 293.2, 298.2 322.7 313.2,

333, 388, 422 349.9, 399.9 353.2, 393.2 313.2, 333.2

353.2, 393.2

313.2, 353.2

P/kPa 156, 225, 570, 1121 1380, 3450, 6900 2000, 5000, 10 000 100, 300 265 200−5000 300 300, 1000 2000

DEEA, diethylethanolamine; DMAEOE, 2-[2-(dimethylamino)ethoxy]ethanol).

corrosiveness toward carbon steel of the conventional alkanolamines is primary > secondary > tertiary.1 For all these reasons, amine blends of tertiary amines (methyldiethanolamine, MDEA) with primary, secondary, and sterically hindered amines (monoethanolamine, MEA; diethanolamine, DEA; and 2-amino-2-methyl-1-propanol, AMP; respectively) have been widely used in industry. Experimental data for the enthalpies of reaction of these tertiary amines is necessary for a good thermodynamic description of the systems, supporting the predictions of phase equilibria and heat effects outside the conditions where the experimental data are available. Experimental values for the heats of reactions are very rare even though the bottleneck of the industry right now is the high cost of regeneration of the amines in CO2 capture operations. Note also that a very large number of amines are proposed as promising solvents without being screened for an important parameter in the cost of operation, i.e., the amount of the heat required in the regeneration. Some of the earlier work carried out in finding the heats of absorption for aqueous tertiary amines is presented in Table 1. In this work, we report the heats of absorption of three structurally similar tertiary amines, N-methyldiethanolamine (MDEA), 3-dimethylamino-1-propanol (3DMAP), and 1dimethylamino-2-propanol (1DMAP), using a flow calorimeter at 298.15, 323.15, and 343.15 K. These last two amines were found by this group (Kadiwala et al.6) to have higher rate constants compared to MDEA (K2 = 0.21 m3·mol−1·s−1 for 1DMAP; K2 = 0.32 m3·mol −1·s−1 for 3DMAP; K2 = 0.12 m3· mol−1·s−1 for MDEA). Heats of absorption of aqueous solution of MDEA (30 wt %) were measured at 322.5 K for different pressures (1.02, 5.14, and 10 MPa). When compared with those published by Acris et al.,7 they have absolute average deviations of 7%. The heat of absorption data for the 30 wt % aqueous MDEA were also compared with other available literature data from different calorimeters, and estimated enthalpies from vapor−liquid equilibrium (VLE) data. Heat of absorption data for a 5 wt % aqueous 3DMAP were compared with 26.3 wt % aqueous 3DMAP data published before us by Rodier et al.8 In addition, the heat of absorption data for aqueous 1DMAP (5 and 30 wt %) were measured and compared with those for MDEA and 3DMAP to understand the structural dependence of heat of absorption values for these tertiary amines. Figure 1 shows the molecular structures of the tertiary amines considered. A correlation model similar to the work done by Merkley9 was developed in MATLAB to predict the temperature dependent equilibrium constants, and the heat of reaction for amine protonation under the studied experimental conditions. The model prediction was compared with the

Figure 1. Molecular structures of the studied tertiary amines.

estimated integral heat of absorption from VLE data. Predicted equilibrium partial pressures and speciation from the model are also reported.

2. EXPERIMENTAL SECTION The equipment used in this study is similar to the one used by Acris et al.,10 and a similar experimental procedure was followed in this work. The heat of mixing of CO2 in aqueous solutions of amines at constant temperature and pressure was measured as a function of the CO2 loading α (molCO2/ molamine), defined as the total molar flow of CO2 divided by the 4954

dx.doi.org/10.1021/ie4041324 | Ind. Eng. Chem. Res. 2014, 53, 4953−4965


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Figure 2. Schematic diagram of the flow calorimeter loop.

total molar flow of aqueous amine. A new mixing flow unit was developed in collaboration with Setaram and then manufactured by Setaram Instrument (Caluire, France). The first unit was first tested in our laboratory for measuring the heat of mixing and then the heat of reaction. The unit was integrated into a C-80 heat conduction differential calorimeter and used to measure the heat produced during the absorption of CO2. Figure 2 describes the overall experimental arrangement. The two fluids, CO2 and aqueous amine solution, enter the mixing unit fed by two ISCO (Model 100 DM) high pressure pumps. The syringe pumps maintain a stable flow rate with a relative uncertainty of 0.2%.The flow rates used varied between 0.01 and 2 mL·min−1 depending upon the molar volume of fluids at the experimental conditions. The volumetric masses of CO2 were taken from the ALLPROPS software.11 The densities of aqueous amine solutions were measured at atmospheric pressure using an Anton Parr densimeter DMA 4500 and validated with available literature data. The uncertainty in density measurements was found to be 5 × 10−3 g·cm−3 with respect to the experimental temperature and pressure, which is acceptable to determine accurately the required flow rates for the aqueous solution. The flow line of the experimental setup is made out of stainless steel tubes of 1.6 mm o.d. and 1.0 mm i.d. The pressure of the system is maintained constant to 0.05 MPa using a Swagelok back-pressure regulator (SS medium to high pressure, KPB1P0D412P2000) placed at the end of the flow line. The pressure is measured by three electronic Keller pressure gauges (LEO1/300 bar/81002C) as shown in Figure 2

with an accuracy of 0.25%. The experiments were carried out at constant temperature, and the temperatures of the entering fluids were controlled by preheaters (±0.05 K) and a Julabo Model F25 water bath (±1 K).The enthalpy of solution (Hs) is calculated from eq 1 using the thermopile signals SM and SBL (mW) and the molar flow rate n (mol·s−1) of the gas (Hs/kJ· molCO2) or the molar flow rate of the amine (Hs/kJ·molamine): Hs =

SM − SBL n

(1)

where SM represents the thermopile signal during the mixing process and SBL is the baseline signal recorded when only the aqueous phase is flowing through the calorimeter. The thermopile sensitivity and temperature coefficients were calibrated using a reference system (ethanol + water) with the enthalpy data provided by Ott et al.12 The Calisto software provided by Setaram records the instantaneous signal values during the experiment. The error on the heat power is related to the accuracy of the thermopile sensitivity of the calorimeter and is estimated (when compared to the literature) to be 2%.

3. MODELING The principle reactions occurring in the CO2−alkanolamine− water system can be obtained from the literature. The main reactions involved in tertiary amines are the following. 4955



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Below the Saturation Point. Below the saturation point, all CO2 is absorbed, and consequently the equilibrium for reaction 6 is not considered. The equilibrium constants are expressed as

ionization of water: K1

H 2O ↔ H+ + OH−

{K1 , ΔH1}

(2)

protonation of tertiary amine:

K1 =

K2

Amine + H+ ↔ AmineH+

{K 2 , ΔH2}

K2 =

K3

{K3 , ΔH3}

+

H + CO3

2−

{K4 , ΔH4}

K3 =

CO2(g) ←→ ⎯ CO2(l)

{HCO2 , ΔH5}

K4 =

∑ ΔniΔHi

(8)

Δn2 = [Amine]f − [Amine]i

(9)

(11)

Initial Concentrations. The molal concentrations can be calculated from stoichiometric equilibria, the charge balance, and component mole balances. The calculation of the initial composition involves only reactions 2 and 3, due to the absence of CO2. The equilibrium constants for reactions 2 and 3 are

K2 =

γH+[H+]i γOH−[OH−]i a H 2O

where γ denotes the activity coefficient and a indicates the activity. The equations to be considered for calculating the initial concentrations are (14)

[Amine]0 = [Amine]i + [AmineH+]i

(15)

[nCO2]0 = [CO2 ]f + [HCO3−]f + [CO32 −]f

(22)

∞ (p − ps ) ⎤ ⎡ υCO (γCO [CO2 ]f ) exp⎣⎢ 2 RT w ⎦⎥ 2

PCO2φCO

(23)

where HCO2 corresponds to the Henry’s law coefficient (mol· kgH2O−1·atm−1) for the absorption of CO2 in pure water and PCO2 (atm) is the equilibrium partial pressure of CO2 over the aqueous solution. Equation 23 disregards any gas phase nonidealities through the use of pressure rather than fugacity. [nCO2]f is the final unreacted CO2 in the solution. It was calculated by subtracting the ionic products from the initial CO2 concentration in the solution [nCO2]i. Along with eqs 20 and 21, the CO2 balance is as follows:

(13)

[AmineH+]i + [H+]i = [OH]i

(21)

2

γH+[H+]i γAmine[Amine]i γAmineH+[AmineH ]i

(20)

[Amine]0 = [Amine]f + [AmineH+]f

HCO2 =

(12)

+

(19)

The simultaneous solution of eqs 16−22 for known values of the equilibrium constants, activity coefficients, and original concentrations of amine and gaseous CO2 yields the final compositions ([Amine]f, [AmineH+]f, [OH−]f, [HCO3−]f, [CO32−]f, [CO2]f). Combining the final and initial solution compositions yields the values of Δni for eqs 8−11 that occur in solution upon mixing. Beyond the Saturation Point. To calculate the final composition for CO2 beyond the saturation loading point, all eqs 2−5 have to be considered. CO2 equilibrium can be defined using the following expression:13

(10)

K1 =

3

γHCO −[HCO3−]f

= [OH−]f + [HCO−3 ]f + 2[CO32 −]f

Δn3 = [HCO3−]f − [HCO3−]i + [CO32 −]f − [CO32 −]i

Δn4 = [nCO2]i − [nCO2]f

γH+[H+]f γCO 2−[CO32 −]f

[AmineH+]f + [H+]f

The change in the number of moles from the initial condition to the final condition (Δni) for each chemical reaction can be expressed as follows (mol/kgH2O): Δn1 = [OH−]f − [OH−]i

(18)

Other equations needed to solve for the final compositions are the charge balance, the amine balance, and the CO2 component balance. They are given as follows:

(7)

i

3

γCO [CO2 ]f a H2O

3

(6)

(17)

γH+[H+]f γHCO −[HCO3−]f 2

The overall heat of solution (Hs) is the sum of the heat of contributions (ΔniΔHi) from each of the five steps: Hs =

γAmineH+[AmineH+]f

(5)

absorption of gaseous CO2 into the solution: HCO2

(16)

γH+[H+]f γAmine[Amine]f

(4)

second ionization of carbonic acid: K4 HCO3− ↔

a H 2O

(3)

first ionization of carbonic acid: H 2CO3 ↔ H+ + HCO3−

γH+[H+]f γOH−[OH−]f

[Amine]0 is the actual stochiometric concentration of the amine (mol/kgH2O). The simultaneous solution of these equations yields the initial composition of the solution ([H+]i, [OH−]i, [Amine]i, and [AmineH+]i). Final Concentrations. The final composition for the aqueous solution after CO2 absorption was calculated for each value of CO2 loading by considering the level of saturation at the loading point.

[nCO2]0 = [nCO2]f + [CO2 ]f + [HCO3−]f + [CO32 −]f (24)

For the known values of K, γ, and the original concentrations of amine and CO2, the final compositions ([nCO2]f, [Amine]f, [AmineH+]f, [H+]f, [OH−]f, [HCO3−]f, [CO32−]f, [CO2]f) for loadings beyond the saturation loading point were calculated by solving these equations. The values of Δni were calculated for 4956



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Table 2. Equilibrium Constants for Different Reactions Used in This Work9 − ln K = A + BT +

C E + D ln T + 2 T T

reaction

A

B

C

D

E

2 4 5 6

−3361.10 3653.20 −213.41 208.02

−0.52 0.60 −0.02 0.03

179 690 −196 970 8 025.10 −7 279.70

534.52 −577.03 37.30 −32.60

−8 891 700 11 763 000 418 330 −316 300

Figure 3. Flow sheet for MATLAB program.

each of the five reactions. To find if the saturation loading point has been reached, first the final composition of the solution is calculated assuming that the saturation loading point has not been reached. From this final composition, the equilibrium partial pressure of CO2 (PCO2) is determined by rearrangement of eq 23. The equilibrium partial pressures of tertiary amine and H2O are determined using the modified extended Raoult’s law:

pi = γixipi *

xi, γi, and pi* are the mole fraction, the activity coefficient, and the vapor pressure of the pure component at the temperature studied. The estimated total pressure is the sum of the partial pressures of CO2, amine, and water. If this total pressure is greater than the actual system pressure, then the saturation loading point has been exceeded and the final composition is recalculated according to eqs 16−21, 23, and 24, assuming gaseous CO2 to be present. This loading was also confirmed experimentally.

(25) 4957



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To calculate the composition and the enthalpy of solution with the model, information about K and ΔH as a function of the temperature for each of the chemical reactions is needed. Correlations for the activity coefficients are also required for calculating the molal concentrations. The equilibrium constants for reactions 2, 4, 5, and 69 are given in Table 2. The temperature dependency of the equilibrium constants is given by the following expression: −ln K = A + BT +

C E + D ln T + 2 T T

U (K 2 , ΔH2) =

j

(26)

∂U (K 2 , ΔH2) =0 ∂ΔH2

(32)

∂U (K 2 , ΔH2) =0 ∂K 2

(33)

Since eq 32 is explicit and linear in ΔH2, the partial derivative given by eq 33 can be calculated explicitly and solved for ΔH2:

ΔH2 =

2

∑j ∑j

(27)

Q c, jΔn2, j N (Δn2, j)2 N

(34)

U(K2, ΔH2) is implicit in K2, and eq 34 cannot be solved explicitly for the equilibrium constant. Therefore, a trial-anderror procedure was followed, and is detailed in Figure 3.

For reactions 2, 4, and 5, a correlation for the activity coefficient ratio γH+γHCO − γ +γ − 3 K1, γ = H OH ; K3, γ = γCO a H2O a H 2O

4. RESULTS AND DISCUSSION MDEA + H2O System. In order to predict the uncertainty of the equipment, the heats of absorption for aqueous 30 wt % MDEA were measured at 322.5 K at 1.02 and 5.14 MPa. Additionally, a few data points were collected at 10 MPa to study the change in the heat of absorption values with respect to pressure. The experimental values are listed in Table S1 (Supporting Information). The experimental values are compared to those published by Acris et al.7 and plotted in Figure 4. As the loading increases, the heat of absorption (kJ·

2

as a function of temperature (273−573 K) and solution ionic strength (I = 5 m) was used (Appendix). For reaction 5, the absorption of gaseous CO2 in the solution, a correlation for the activity coefficient (K5,γ) was taken from Patterson et al.14 The first step in the modeling procedure, using the measured overall enthalpies of solution and literature values of the equilibrium constants and heat of reactions for chemical reactions 2, 4, 5, and 6 along with activity coefficient correlations, is to determine a thermodynamically consistent set of K and ΔH values for the amine protonation reaction (3).The procedure for finding the best values of K and ΔH is given in Figure 3. The relationship between the measured heat and the equilibrium constants and the enthalpy changes is given by eq 7. Δni values are calculated by eqs 8−11, and then the energy released by the amines protonation reaction alone (Qc,j) for which reaction both K and ΔH are unknown is calculated by other reactions for which K and ΔH are known from the total measured heat: Q c, j = His, j −

(31)

The best values of K2 and ΔH2 for a given temperature are those that minimize U(K2, ΔH2), which satisfy the following criteria:

The model was used to find the values of K and ΔH for the amine protonation reaction (3) by fitting the experimentally measured values and the calculated values of the total heat of solution. The heat of reaction was derived with a multiple linear regression using the van’t Hoff equation as ⎛ ∂ ln K E⎞ = ΔH = R ⎜C − DT − BT 2 + 2 ⎟ RT ⎝ ∂T T⎠

∑ uj

∑ Δni ,jΔHi (28)

i≠2

where the subscript j refers to the loading at a single point. This corrected heat value is calculated for each data point. Because this corrected heat represents the net effect of the amine protonation reaction alone, Qc,j may be expressed as follows: Q c, j = ΔH2Δn2, j

(29)

Figure 4. Heat of absorption of CO2 in aqueous 30 wt % MDEA at 322.5 K and different pressures.

The best values of K2 and ΔH2 at each temperature were calculated by a least-squares analysis of eq 29. The normalized error square (uj) for a single data point is given as

molCO2−1) remains constant up to the saturation loading point and then starts decreasing beyond that point. A physical interpretation of these trends is that, prior to the saturation loading point, all CO2 fed to the calorimeter is completely absorbed and the total heat released upon the absorption is directly proportional to the amount of CO2. Therefore, as the loading increases the values of the numerator in ratio kJ· molCO2−1 between total heat released (kJ) and number of moles

2

uj =

(Q c, j − ΔH2Δn2, j) N

(30)

where N refers to the number of data points and K2 and ΔH2 are functions of temperature only. The normalized error square sum (U) is the summation of the normalized square error (uj) over all the loading values: 4958



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reactions well compared to the physical dissolution beyond the saturation. The method of finding the saturation loading point can be explained by plotting kJ·molMDEA−1 versus the loading (Figure S1 in the Supporting Information). The plot has a positively sloped straight line from zero loading until the saturation loading point, and after the saturation loading point, for low pressure the data remain constant, but for the high pressure the data show a negative slope. This may be due to the evaporation of water and MDEA into the vapor phase causing the total heat offset by the endothermic heat effect. Therefore, the heat of absorption decreases by increasing the loading after that saturation loading point at high pressures. The saturation loadings were found to be 0.95, 0.98, and 1.01 at 1.02, 5.14, and 10 MPa. Acris et al.7 reported somewhat higher values of 1.04 and 1.2 at 1.02 and 5.14 MPa. The heat of absorption values for CO2 in aqueous 30 wt % MDEA were measured with different calorimetric techniques such as an isoperibol calorimeter, a flow calorimeter, and a semidifferential calorimeter. The heat of absorption values at infinite dilution determined by different calorimeters were compared with the data obtained in this work. Carson et al.29 used an isoperibol calorimeter and reported the value as −48.7 kJ·mol−1 at 0.27 MPa. Acris et al.7 employed a flow calorimeter and reported −59.2 kJ·mol−1 (0.52 MPa), −57.1 kJ·mol−1 (1.02 MPa), and −56.8 kJ·mol−1 (5.14 MPa) at 322.5 K. Kim et al.30 reported −59.17 and −57.7 kJ· mol−1 at 313.15 K using a semidifferential calorimeter at 0.3 and 1 MPa. The infinite dilution values obtained in this work are −56.1 kJ·mol−1 (1.02 MPa), −57.2 kJ·mol−1 (5.14 MPa), and −57.8 kJ·mol−1 (10 MPa) at 322.5 K. Note that, with a 2% accuracy, these values are considered constant, and the influence of pressure in this range for this work and others can be considered as negligible. The model was used to find the heat of the protonation reaction, when tertiary amines react with CO2. The values estimated for 30 wt % aqueous MDEA in this work are −52.4 kJ·mol−1 (322.5 K) at 1.02 MPa and −47.6 kJ·mol−1 (322.5 K) at 5.14 MPa. 3DMAP + H2O System. The heat of absorption values for 3-dimethylamino-1-propanol (5 wt %) were obtained at 298.15, 323.15, and 343.15 K for 2 MPa. The experimental values are reported in Table S2 (Supporting Information). The trend in the heat of absorption values (kJ·molCO2−1) with respect to loading (molCO2/mol3DMAP) is shown in Figure 5. Rodier et al.8 measured the heat of absorption values for 26.4 wt % 3DMAP at 313.15 and 353.15 K, and those values are also plotted in Figure 5 to show the temperature dependency of the heat of absorption. The saturation loading points were found to be 1.27, 1.06, and 0.94 molCO2/mol3DMAP at 298.15, 323.15, and 343.15 K and are shown in Figure 6. Rodier et al.8 reported the saturation limit for 3DMAP at 313.15 K as 1.38 molCO2/ mol3DMAP and at 353.15 K as 1.06 molCO2/mol3DMAP, in good agreement with this work. As the temperature increases, the saturation loading point decreases. Comparing the results of the heat of absorption values with Rodier et al.,8 it was observed that the change in the amine concentration did not influence much the values in contrast with the influence of temperature. In agreement with their work, the limits of solubility are superior to unity at lower temperatures. Both chemical reactions and physical dissolution contribute to this absorption capacity. Rodier et al.8 explained the high solubility observed in 3DMAP by the alkalinity (pKa ≈

of CO2 (molCO2) remain constant until the saturation loading. Beyond the saturation loading, the absorption of CO2 is not proportional to the total heat released due to saturation and the total heat measured is the heat required just to saturate the solution. Thus, beyond the saturation loading point, the value of the numerator (kJ) decreases proportionally due to an increase in the amount of CO2. The integral heats of absorption values obtained in this work were compared to the predicted heat absorption values from the vapor−liquid equilibrium (VLE) data obtained from the literature.15−25 The solubility data for CO2 in aqueous MDEA were used, and the differential heat of absorption values were obtained using the Gibbs− Helmholtz equation: ⎡ ⎤ ΔHdiff ⎢ ∂ ln PCO2 ⎥ ⎢ ⎥ =− R 1 ⎣ ∂ T ⎦α

()

(35)

In order to be compared to the integral heat of absorption, the differential heats of absorption were fitted with a modified Gompertz equation:26 −ΔHdiff = a − bc d

−α

(36)

where a = 62.7824, b = 427.3773, c = 0.0053, and d = 1.8648. The standard error of fitting is ±6.4 kJ·mol−1. Then, eq 36 was integrated with respect to the loading, α: −ΔHint = −

1 α

∫0

α

⎡ bEi[d α ln c] bE [ln c] ⎤ + i ⎢aα − ⎥ ⎣ ln d ln d ⎦ (37)

Integration of eq 37 was done using the Mathematica program for different loadings. Figure 4 shows the trend of the predicted integral enthalpy which is very much similar to the experimental integral enthalpy in below saturation loading and deviating more as it approaches the saturation loading point. Using the binary parameters for an electrolyte nonrandom two liquid (ENRTL) model suggested by Blanchon le Bouhelec et al.,27 the differential heat of absorption values were predicted using a flash calculation and were then converted into integral heat of absorption. The reactive flash calculations were done using the work of Blanchon le Bouhelec et al.27 The flash calculation used in this work is similar to the procedure described by Hilliard.28 The fitted values for the modified Gompertz equation were a = 49.2229, b = 55.0673, c = 2.1895 × 10−16, and d = 41.2015. The standard error of fitting is ±1.5 kJ·mol−1. The predicted values from the ENRTL model below the saturation loading point were lower than the experimental integral heat of absorption and the predicted values from VLE data using the Gibbs−Helmholtz equation. However, the values beyond the saturation loading point are closer to the experimental integral heat of absorption values. This may be due to the data used in the VLE and ENRTL model. The ENRTL model prediction for solubility uses the correlation for the equilibrium constant related to the amine protonation. The choice of the correlation leads to large deviations in enthalpies of protonation, and 80% of the enthalpy of solution was contributed by the amine protonation as reported by Acris et al.13 From this work, it can be concluded that the ENRTL model used represents the physical dissolution very well compare to chemical reactions. Whereas, modified Gompertz equation from the VLE experiments represent the chemical 4959



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Figure 5. Heat of absorption of CO2 in aqueous 3DMAP at different temperatures and 2 MPa.

Figure 7. Equilibrium partial pressure of CO2 predicted by the model in aqueous 3DMAP (5 wt %) at different temperatures.

Figure 6. Heat of absorption of CO2 in aqueous 3DMAP at 2 MPa and different temperatures. Line intersections, for each temperature, were used to determine the saturation loading.

Figure 8. Equilibrium concentrations of CO2 loaded aqueous 3DMAP (5 wt %) predicted by the model at 323.15 K.

amines. The positive slope of the trend of these ions up to the saturation point indicates the occurrence of a chemical reaction. After the saturation loading point, these ions showed a linear trend, and the CO2 concentration showed an exponential rise. 1DMAP + H2O System. In order to study the structural dependency of the heat of absorption of CO2 in amines, 1DMAP (an isomer of 3DMAP) was chosen and the heat of reaction for the aqueous amine solutions (5 and 30 wt %) were measured at 298.15, 323.15 and 343.15 K at 2 and 5 MPa. The experimental values are presented in Tables S3−S5 (Supporting Information). The trend in the heat of absorption values with respect to loadings is given in Figure 9. Increases in the concentrations of amines did not affect the heat of absorption values, but an increase in pressure decreased the values similar to the decrease in temperature. The heat of absorption value below the saturation for aqueous 1DMAP was higher than that of 3DMAP. Comparing the dissociation constants, 3DMAP with pKa ≈ 9.54 at 298.15 K6 and 1DMAP with pKa ≈ 9.50 at 298.15 K6 have similar alkalinities; therefore, the alkalinity is not the reason for this high heat of absorption in 1DMAP. Comparing the chemical reaction rate constants (K2) of 1DMAP with 0.21 m3·mol−1·s−1 and for 3DMAP with 0.32 m3·mol−1·s−1,6 it can be concluded that 3DMAP reacts faster

9.51 at 298.15 K). The experimental data of this work and Rodier et al.8 were correlated to find the temperature dependent equilibrium constants and the heats of reactions for the protonation reaction of 3DMAP are calculated. They can be represented as follows: ⎛1⎞ ln K 2 = −10063⎜ ⎟ + 11.15 ⎝T ⎠ ⎛1⎞ ΔH2 = 0.064⎜ ⎟ + 28.71 ⎝T ⎠

(R2 = 0.99) (R2 = 0.99)

(38)

(39)

The percentile absolute average deviations (AAD%) of the thermodynamic consistent ln K2 and ΔH2 are 6.4, 7.3, 2.6, 3.6, and 6.0 for 298.15, 313.15, 323.15, 343.15 and 353.15 K, respectively. The predicted partial pressures for all temperatures are presented in Figure 7. The predicted molalities of the species are given in Figure 8. Note that an increase in loading increases the formation of [HCO3−] and [3DMAPH+] ions. This shows that the protonation reaction is the main contributor for both the heat of reaction and the solubility in aqueous tertiary 4960



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Figure 10. Heat of absorption of CO2 in aqueous 1DMAP for 2 and 5 MPa and different temperatures. Line intersections, for each temperature, were used to determine the saturation loading.

Figure 9. Heat of absorption of CO2 in aqueous 1DMAP at 298.15, 323.15 and 343.15 K and at 2 and 5 MPa.

with CO2 than 1DMAP. Activation energies for these tertiary amines were estimated using a base catalysis mechanism by Kadiwala et al.6 and showed that 3DMAP (75.55 kJ·mol−1) has a higher activation energy than 1DMAP (62.55 kJ·mol−1); therefore, the energy needed for the reaction to occur in 1DMAP is less than in 3DMAP and the additional heat energy needed to absorb CO2 in 1DMAP over 3DMAP should come from another source. The difference in the molecular structures of these amines is mainly the position of −OH groups in the molecule as shown in Figure 1. The standard state enthalpy change of 2-amino-2-methyl-1-propanol (AMP), one of the primary sterically hindered amines, has the highest standard state enthalpy change of conventional amines.31 A similar increase in enthalpy was observed by comparing 2-piperidineethanol (53.8 kJ·mol−1) and 1-piperidineethanol (40.9 kJ· mol−1).31 The negative values of excess heat of mixing with water were observed to be quite different also for the two amines and were higher, in absolute values, for 3DMAP than for 1DMAP (by ∼700 J/mol) in a recent work.32 Nezamloo33 suggests that the difference in the heat of mixing between the two amines is mainly due to a weaker hydrogen bonding between the nitrogen atom and the interacting water molecule in 1DMAP due to a steric effect obtained by the position of the OH. This could also explain the lower kinetic rate of CO2 in 1DMAP, mentioned earlier, and could also explain the increase in the heat of reaction when compared to 3DMAP. A VLE study of these two aqueous amine systems should shed more light on this behavior. The saturation loading points determined for 30 wt % 1DMAP at 2MPa for 298.15, 323.15, and 343.15 K were found to be 0.75, 0.71, and 0.69, which are lower than those for 3DMAP at these temperatures (as shown in Figure 10). This may be due to the position of the −OH group from the amine group and the higher electron withdrawing effect of the −OH group than of the amine group. Xu et al.31 compared the pKa values of morpholine (MO) and piperidine (PD). They explained that the lone pair of electrons in the nitrogen atom in the MO ring was attracted by the oxygen atom so that the electronegativity of morpholine increases and its ability to combine with H+ ion in water decreases. Since CO2 is an acidic gas and 1DMAP is a tertiary amine (no carbamate formation), the behavior of the system will be reduced mainly to a protonation reaction. This could be a reason for the lower

loading capacity for 1DMAP compared to 3DMAP. The experimental data of this work were correlated to find the temperature dependent equilibrium constants and the heats of reactions for the protonation reaction of 3DMAP are calculated. They are represented as follows: ⎛1⎞ ln K 2 = −6774⎜ ⎟ − 5.126 ⎝T ⎠

⎛1⎞ ΔH2 = 27456⎜ ⎟ − 40.87 ⎝T ⎠

(R2 = 0.99)

(R2 = 0.99)

(40)

(41)

Percentile absolute average deviations (AAD%) of the thermodynamically consistent ln K2 and ΔH2 for 5 wt % 1DMAP at 2 MPa are 3.1, 4.4, and 4.7 at 298.15, 323.15, and 343.15 K. For 30 wt % 1DMAP at 2 MPa, the calculated AAD% are 9.7, 9.5, and 5.5 at 298.15, 323.15, and 343.15 K. For 30 wt % 1DMAP at 5 MPa, the calculated AAD% are 5.3, 6.4, and 6.5 at 298.15, 323.15, and 343.15 K. The predicted partial pressures for all temperatures are given in Figure 11. The predicted molalities of the species are given in Figure 12. From Figure 12, it can be noted that the increase in

Figure 11. Equilibrium partial pressure of CO2 predicted by the model in aqueous 3DMAP (5 and 30 wt %) at different temperatures (298.15, 323.15, and 343.15 K) and pressures (2 and 5 MPa). 4961



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those for 3DMAP (this work and Rodier et al.8) and 2-[2(dimethylamino)ethoxy)]ethanol (DMAEOE; Rodier et al.8) in Figure 13. It can be observed that as the steric hindrance of the

Figure 12. Equilibrium concentrations of CO2 loaded aqueous 1DMAP (5 wt %) predicted by the model at 323.15 K.

loading increases the formation of ions HCO 3 − and 1DMAPH+. This shows that the protonation reaction and formation of carbonate ions (second ionization) are the main contributors for the heat of reaction and solubility in aqueous 1DMAP.The first ionization is negligible until the saturation loading point is reached and increases rapidly after the saturation point. Comparing 3DMAP and 1DMAP speciation, it is observed that the rate formation reaction products are dissimilar at different loadings. The maximum concentration of CO3− ions occurs near the saturation point of 1DMAP and shows that the deprotonation of HCO3− ions is faster in this region, and this may be a reason for the higher energy contribution of 1DMAP compared to 3DMAP. The physical absorption region is reached faster in 1DMAP than in 3DMAP. Comparison of Aqueous Tertiary Amines. In order to compare the heat of absorption values of 3DMAP, which has a higher loading capacity and higher reaction rate than MDEA and triethanolamine (TEA), the integral heat of absorption values of 30 wt % MDEA and TEA were calculated from the literature. The calculated integral heats of absorption for 30 wt % MDEA were reported previously. VLE data for TEA were taken from Jou et al.34 and Chung et al.35 and correlated with the ENRTL model to get the binary interaction parameters. The heat of absorption values were derived using the Gibbs− Helmholtz equation in VLE data using a flash calculation at 313.15 K.28 The differential enthalpies were fitted with the modified Gompertz equation. The fitted equation is given as

Figure 13. Comparison of the heat of absorption values of (≈30 wt %) aqueous tertiary amines at 313.15 K.

amino group is increased by the alkyl groups the heat of absorption decreases. The order of the heat of absorption values is TEA < MDEA < 3DMAP < DMAEOE < 1DMAP. As the number of −OH group increases, the heat of absorption values decrease. The position of the −OH group from the amino also influences the heat of absorption as can be observed by comparing the heat of absorption values of 3DMAP and DMAEOE. Kadiwala et al.6 observed that 3DMAP has higher CO2 loading capacity and reaction rate than other tertiary amines such as MDEA and TEA. Since tertiary amines are also less environmentally friendly because of their low biodegradability,36 the biodegradability of 3DMAP has to be evaluated. In view of the recent trends to blending primary and secondary amines with tertiary amines, 3DMAP should be a better candidate compared to MDEA and therefore of interest for CO2 capture.

5. CONCLUSION A flow calorimeter was used to measure the heat of absorption of CO2 in three tertiary amines (MDEA, 3DMAP, and 1DMAP) at different temperatures and pressures. The increase in pressure did not affect the heat of absorption values for aqueous MDEA, but increased the saturation loading point. The heat of absorption values of structurally similar tertiary amines (1DMAP and 3DMAP) compared to MDEA were measured at 298.15, 323.15, and 343.15 K at 2 MPa. The measured heat of absorption values of aqueous 3DMAP (5 wt %) were compared with the values of Rodier et al.8 (26.3 wt % 3DMAP) measured at 313.15 and 353.15 K. An increase in the concentration of 3DMAP did not change much the heat of absorption values compared to the effect of temperature, but the CO2 loading capacity increased. Using the experimental data (this work and Rodier et al. 8 ) and a rigorous thermodynamic model, the temperature dependent equilibrium constant for the protonation reaction was regressed. The experimental data and the correlated model data for the heat of absorption values was within 6 AAD%. The model was used to

−α

−ΔHdiff = 50.01 − 26.25(4 × 10−8)33.82

(42)

The standard error in the fitted equation is ±1.42 kJ·mol−1. Then eq 42 was integrated with respect to loading, α: −ΔHint = −

1 α

∫0

α

⎡ ⎢50.01α ⎣

26.25Ei[33.82α ln 4 × 10−8] ln 33.82 26.25Ei[ln 4 × 10−8] ⎤ ⎥ + ln 33.82 ⎦

−

(43)

where Ei is the Gompertz constant (0.596 347). The integral heat of absorption values were calculated and compared with 4962



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predict the partial pressures and speciation of the aqueous solutions and interpret the experimental heat of absorption values and the vapor−liquid behavior of the tertiary amine systems. The heat of absorption values of 1DMAP (structurally similar to 3DMAP) were measured at 298.15, 323.15, and 343.15 K at 2 and 5 MPa for 5 and 30 wt % aqueous solutions. It was observed the heats of absorption of CO2 in 1DMAP were larger than in 3DMAP, which is mainly due to the position of −OH group in the molecule. After regression using the model, it was observed that the concentration of CO32− is higher in 1DMAP when compared to 3DMAP at lower loadings, indicating that the second ionization reaction was faster. The saturation loading of 1DMAP (0.75 molCO2/mol1DMAP) was very much less than that of 3DMAP (1.27 molCO2/mol3DMAP). The heat of absorption values of 3DMAP were compared with those of well-known tertiary amines (MDEA, TEA, DMAEOE), and were found to be similar. However, 3DMAP had more loading capacity for CO2 and reacted faster with it. It will be therefore beneficial to consider 3DMAP as a possible replacement to MDEA in CO2 removal operations after testing it for degradation and corrosion.

C = 4.86989 × 10−3 − 6.65184 × 10−5t + 1.69534 × 10−7t 2

t is in °C in eqs A.6 and A.7. ⎛ α ⎞ DH(βI1/2) = ⎜ 3 ⎟[(1 + βI1/2) − 2 ln(1 + βI1/2) ⎝β I⎠ − (1 + βI1/2)−1]

β = 1.64801 − 2.83246 × 10−3t + 4.67896 × 10−6t 2 (t > 107.7°C)

(t > 107.7°C)

⎛ 2.33752 × 104 ⎞3/2 α = 1.17202ρ1/2 ⎜ ⎟ DT ⎝ ⎠

D=

ρ is the density of pure water and was calculated by

(A. 13)

where t is in °C. For reaction 2 γ γ + K 2, γ = Amine H γAmineH+

ln γH+ = f + mB + m2C ⎡ I1/2 ⎤ ⎥ f = −Aϕ⎢ ⎣ 1 + bI1/2 ⎦

S = −2.97627 + 4.80688 × 10−2T − 2.69280 × 10−4T 2 4

B = β (0) − β (1) exp( −αI1/2)

1 − (1 + 2I1/2 − 2I ) exp( −2I1/2) 4I

ϕ = 1 − DH(βI

) + BI + CI

2

(A.4)

(A.17)

(1)

ln γMEAH+ =

(A.5)

−AI1/2 − 0.0750[HCO3−] 1 + bBI1/2

(A.18)

A is the Debye−Hückel limiting slope (as defined in eq A.11), b = 4.5 × 10−8, B = 50.3 × 108(DT)−1/2, and D is the dielectric constant for pure water (A.12).

B = 4.43570 × 10−2 + 2.14240 × 10−4t − 1.21582 × 10−6t 2

(A.16)

β = 0.183 52, β = 0.255 03, α = 2, and C = −0.000 59. For the initial concentration calculation, it was assumed that γMDEAH+ = γMEAH+. The correlation for γMEAH+ was taken from Deshmukh and Mather:44 (0)

ϕ is the osmotic coefficient and given by the following expression:42 1/2

(A. 15)

b = 1.20, I is the ionic strength (as defined in eq A. 2), and Aϕ is the Debye−Hückel limiting slope (as defined in eq A.11).

(A.3)

where T is in kelvin. F (I ) =

(A. 14)

For the initial concentration calculation, the activity coefficient of amine was assumed to be unity and the activity coefficient of H+ is given by43

where [i] is the molality of component i and zi is the ionic charge of component i. S is the Debye−Hückel coefficient and is given by9

+ 5.58004 × 10−13T 5

(A.12) 42

ρ = 1.00157 − 1.56096 × 10−4t − 2.69491 × 10−6t 2

(A. 2)

+ 7.49524 × 10 T − 1.02352 × 10 T

5321 + 233.76 − 0.9297T + 1.417 × 10−3T 2 T − 8.292 × 10−7T 3

In eq A.1, for a = 1, reaction 2, ionization of water, 40 P10 = −0.611 39, P11 = 87.645, P12 = 1.6698 × 10−6, and P13 = −0.244 56. For a = 3, reaction 4, first ionization of carbonic acid,41 P10 = −0.688 379, P11 = 136.722, P12 = 1.935 21 × 10−6, and P13 = −0.124 789. For a = 5, reaction 5, second ionization of carbonic acid,14 P10 = 9.567 41× 10−3, P11 = −47.4028, P12 = 2.302 64 × 10−7, and P13 = −0.159 489.

−9

(A.11)

D is the dielectric constant of pure water and calculated as42

(A.1)

−7 3

(A. 10)

α is the Debye−Hückel limiting slope and is defined as42

⎡ ⎤ P 2S I + ⎢P10 + 11 + P12T 2 + P13F(I )⎥I ⎣ ⎦ T 1+ I

I = 0.5 ∑ [i]zi

(A.9)

β = 0.980075 − 7.07654 × 10−3t + 2.97433 × 10−5t 2

APPENDIX Equation A.1 provides the correlation for the activity coefficient, with values of a and Pi for reactions 2, 4, and 5:

− 0.157ϕ

(A.8)

β is the ion-size parameter and is defined as

■

−log Ka, γ =

(A.7)

(A.6) 4963



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For final concentrations, the activity coefficients for amine (γAmine) and protonated amine (γAmineH+) were calculated using the suggestion by Edwards et al.:45 ln γi =

−2.303A γ zi2I1/2 1 + BaI1/2

pi = vapor pressure of component i PCO2 = partial pressure of CO2 (MPa) Qc = corrected heat (kJ·kgH2O−1) R = universal gas constant (8.314 kJ·kmol−1·K−1) ΔHabs = heat of absorption (kJ·molCO2−1) ΔHi = heat of reaction reaction i (kJ·mol−1)

+ 2 ∑ βi , jmj j

(A.19)

Greek Symbols

where Aγ is the Debye−Hückel limiting slope (0.509 at 298.15 K in water), B is a function of the temperature and the dielectric constant of the solvent (water), and I is the ionic strength as defined in eq A. 2. zi is the charge number and it is zero for neutral molecule species (amine). The quantity a is an adjustable parameter measured in angstrom units which roughly corresponds to the effective size of the hydrated ions, and βi,j represent the net effect of various short-range two-body forces between different molecular and ionic solutes called binary interaction parameters. The binary interaction parameters of all tertiary amines are assumed to be the same as the binary interaction parameters of MDEA, and are listed in Table S6 (Supporting Information). The first term in eq A.19 represents the contribution of the electrostatic forces and the second term is the representation of short-range van der Waals forces. For reaction 6, the correlation for the activity coefficient is given by46 ln K5, γ =

Q1 T

+ Q 2 + Q 3T

Q i = −biI + ci[1 − exp(diI1.3)]

α = loading (molCO2/molamine) γ = activity coefficient Φ = fugacity coefficient Subscripts

■

(A.20) (A.21)

ASSOCIATED CONTENT

S Supporting Information *

Experimental data for heat of absorption values as mentioned in the figures and the binary interaction parameters used in the modeling work. This material is available free of charge via the Internet at http://pubs.acs.org.

■

REFERENCES

(1) Chakma, A.; Tontiwachwuthikul, P. Designer Solvents for Energy Efficient CO2 Separation from Flue Gas Streams. Greenhouse Gas Control Technologies; Eliasson, B., Riemer, P. W.F., Wokaum, A., Eds.; Elsevier Science Ltd.: Amsterdam, The Netherlands, 1999; pp 35−42. (2) Lee, J. I.; Otto, F. D.; Mather, A. E. The Solubility of H2S and CO2 in Aqueous Monoethanolamine Solutions. Can. J. Chem. Eng. 1974, 52, 803. (3) Kim, I.; Svendsen, H. F. Heat of Absorption of Carbon Dioxide (CO 2 ) in Monoethanolamine (MEA) and 2-(Aminoethyl)ethanolamine (AEEA) Solutions. Ind. Eng. Chem. Res. 2007, 46, 5803. (4) Mathonat, C.; Majer, V.; Mather, A. E.; Grolier, J. P. E. Enthalpies of Absorption and Solubility of CO2 in Aqueous Solutions of Methyldiethanolamine. Fluid Phase Equilib. 1997, 140, 171. (5) Rochelle, G. T.; Bishnoi, S.; Chi, S.; Dang, H.; Santos, J. Research Needs for CO2 Capture from Flue Gas by Aqueous Absorption/Stripping; Final Report for P.O. No. DE-AF26-99FT01029; Sept 26, 2000. Texas Carbon Management Program, Publications. http://research.engr. utexas.edu/rochelle (accessed Feburary 2014). (6) Kadiwala, S.; Rayer, A. V.; Henni, A. Kinetics of Carbon dioxide (CO2) with Ethylenediamine, 3-Amino-1-Propanol in Methanol and Ethanol, and with 1-Dimethylamino-2-Propanol and 3-Dimethylamino-1-Propanol in Water Using Stopped-Flow Technique. Chem. Eng. J. 2012, 179, 262. (7) Arcis, H.; Rodier, L.; Ballerat-Busserolles, K.; Coxam, J.-Y. Enthalpy of Solution of CO2 in Aqueous Solutions of Methyldiethanolamine at T = 322.5 K and Pressure up to 5 MPa. J. Chem. Thermodyn. 2008, 40, 1022. (8) Rodier, L.; Ballerat-Busserolles, K.; Coxam, J.-Y. Enthalpy of Absorption and Limit of Solubility of CO2 in Aqueous Solutions of 2Amino-2-Hydroxymethyl-1,3-Propanediol, 2-[2-(Dimethyl-Amino) Ethoxy] Ethanol, and 3-Dimethyl-Amino-1-Propanol at T = (313.15 and 353.15) K and Pressures up to 2 MPa. J. Chem. Thermodyn. 2010, 42, 773. (9) Merkley, K. Enthalpies of Solution of CO2 in Aqueous MDEA Solutions. M.S. Thesis, Brigham Young University, Provo, UT, USA, 1987. (10) 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. (11) Lemmon, E.; Jacobsen, R.; Penoncello, S.; Beyerlein, S. ALLPROPS 4.2Computer Programs for Calculating Thermodynamic Properties of Fluids of Engineering Interest; Center for Applied Thermodynamic Studies, University of Idaho, Moscow, ID, USA, 1995. (12) Ott, J. B.; Cornett, G. V.; Stouffer, C. E.; Woodfield, B. F.; Guanquan, C.; Christensen, J. J. Excess Enthalpies of (Ethanol + Water) at 323.15, 333.15, 348.15, and 373.15 K and from 0.4 to 15 MPa. J. Chem. Thermodyn. 1986, 18, 867. (13) Arcis, H.; Rodier, L.; Ballerat-Busserolles, K.; Coxam, J.-Y. Modeling of (Vapor + Liquid) Equilibrium and Enthalpy of Solution

b1 = −653.3, b2 = 3.9398, b3 = −0.006 491, c1= −155.3, c2 = 1.1695, c3 = −0.001 981, d1 = −2.504, d2 = −2.276, and d3 = −2.163.

■

diff = differential int = integral

AUTHOR INFORMATION

Corresponding Author

*E-mail: [email protected]. Notes

The authors declare no competing financial interest.

■

ACKNOWLEDGMENTS The financial support of the Natural Sciences and Engineering Research Council of Canada (NSERC), Natural Resources Canada (NRCan), Canadian Foundation for Innovation (CFI), Petroleum Technology Research Centre (PTRC), and the Acid Gas Removal Laboratory (AGR Lab, University of Regina) are gratefully acknowledged. The authors thank Dr. Gary Rochelle and Peter Frailie for providing the flash calculation procedure in ASPEN PLUS.

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NOMENCLATURE T = temperature (K) Hs = heat of solution (kJ·molamine−1) Ki = equilibrium constant for reaction i ni = number of moles of component i pKa = dissociation constant 4964



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