Simulation of Enthalpy and Capacity of CO2 Absorption by Aqueous

A model has been developed to predict the CO2 capacity of amine-based solvent systems as well as the enthalpy associated with absorption/desorption...
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2002

Ind. Eng. Chem. Res. 2008, 47, 2002-2009

Simulation of Enthalpy and Capacity of CO2 Absorption by Aqueous Amine Systems Nichola McCann,*,†,‡ Marcel Maeder,‡ and Moetaz Attalla† CSIRO Energy Technology, P. O. Box 330, Newcastle, New South Wales 2300, Australia, and Department of Chemistry, UniVersity of Newcastle, Callaghan, New South Wales 2308, Australia

A model has been developed to predict the CO2 capacity of amine-based solvent systems as well as the enthalpy associated with absorption/desorption. This model can be used to accurately predict the behavior of well-characterized solvent systems under a range of different conditions. Alternatively, the model can be used to estimate the properties of less well-defined systems as part of an initial rapid screening procedure. Investigation into the effects of varying amine basicity and degree of carbamate formation indicates that there is considerable room for improvement on the standard MEA (monoethanolamine) system in terms of both capture capacity and enthalpy of CO2 desorption. Introduction The global issue of climate change is receiving everincreasing attention.1-7 One of the main focuses of this attention is the reduction of greenhouse gases produced by anthropogenic activities, specifically the reduction of CO2 emissions from activities such as power generation.8,9 One technology which has the potential to make a significant impact on emissions is CO2 capture and storage.1,8,10-12 A simplified schematic of the aqueous amine CO2 capture process currently in commercial use is shown in Figure 1. Flue gas (containing ∼12% CO2) is fed into the absorber column at relatively low temperatures. The CO2 is absorbed into the solvent while the CO2-free gas (mostly N2) is released to the atmosphere. The CO2-rich solution is then transferred to the stripper where comparatively high temperatures are used to desorb the CO2 from the solution. The relatively pure CO2 can then be compressed and sent for storage. The CO2-lean solution is returned to the absorber and the cycle begins again.8 Clearly there is considerable scope for improvements to the engineering of the system; however, this is not the topic of this work and will not be discussed further. There is also considerable scope for improvements to the chemistry of the system. The aqueous amine based solvents that have to date received the greatest attention scientifically and commercially for the removal of CO2 have been systems based on monoethanolamine (MEA). This is because MEA has a high reactivity, low solvent cost, low molecular weight (giving a high absorption capacity on a weight basis), reasonable thermal stability, and low thermal degradation rate. MEA is far from being an ideal solvent, however, as it is susceptible to oxidative degradation, forms thermally stable salts in the presence of NOx- and SOx-bearing gases, is corrosive, forms a relatively stable carbamate, and has a high enthalpy of reaction, corresponding to a large solvent regeneration energy requirement.9 This last point is one of the major sources of cost in the MEA system, and relates to the subject of this work. The majority of current investigations into the potential of novel systems for use as CO2 capture agents focus on extensive vapor-liquid equilibrium (VLE) measurements, sometimes * To whom correspondence should be addressed. Tel: +61 2 4960 6125. Fax: +61 2 4960 6111. E-mail: [email protected]. † CSIRO Energy Technology. ‡ University of Newcastle.

Figure 1. Cyclic CO2 capture using aqueous amine solvents.

coupled with enthalpy measurements. The enthalpy of desorption describes the energy uptake by the chemical process of desorption and so gives vital information for the determination of regeneration energy requirements. The solubility of CO2 in known solvent systems has been thoroughly modeled, based on VLE measurements.13-19 However, this work has not been extended to accurately calculate the enthalpy associated with the process. Currently, little information is available on the enthalpy of reaction of CO2 with amines, and as a result it is difficult to incorporate this factor into the design of an efficient capture plant.20 Instead, much of the research reported to date addressing the issue of high regeneration energy (particularly with MEA) has focused on improving plant design.21,22 Nevertheless, significant reduction to the solvent regeneration energy requirements can be achieved by optimization of the CO2 desorption enthalpy. This work describes the development of a model to describe the equilibria involved in CO2 absorption. Given a thorough understanding of the equilibria, the enthalpy associated with the reaction is then modeled. The model is tested against experimentally obtained results for an aqueous solution of MEA to confirm the validity of the model (with respect to both solubilities and enthalpies). At this point individual species contributions to both solubility and enthalpies are calculated. Second, the effect of the amine basicity and carbamate stability are investigated, allowing insight into the possible potential of amine-based solvents in CO2 capture. Finally, the model is adapted for application to novel solvent systems. The modified version is also tested against a known solvent system and found to be adequate.

10.1021/ie070619a CCC: $40.75 © 2008 American Chemical Society Published on Web 02/23/2008

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very small.23 Consequently, eqs 1 and 2 are typically combined as shown in eq 3.23

Scheme 1

CO2(g) h CO2(aq)*

(3)

where CO2(aq)* represents both dissolved CO2 (CO2(aq)) and hydrated CO2 (H2CO3). This representation will be used in this work. 2. Formation of Bicarbonate. The following equilibria describe the deprotonation of CO2(aq)*: In summary, this work describes a powerful tool for initial screening of potential new capture solvents. This tool can also be used to provide valuable information to assist in directing the development of new solvent systems. Modeling

CO2(aq)* h HCO3- + H+

(4)

HCO3- h CO32- + H+

(5)

3. Formation of Carbamate.

Equilibria. The reaction under investigation involves passing gaseous CO2 (diluted with N2) through an aqueous solution of MEA until equilibrium is reached. The reaction of CO2 with aqueous amines (A) can theoretically be represented in three individual steps as shown in Scheme 1: 1. Dissolution of the gaseous CO2. This step is a purely physical process and must occur prior to further reaction of the CO2. Under absorber conditions, just small quantities of CO2 remain in the form CO2(aq), and for speciation studies (discussed in the Model Validation section), this step will be grouped with bicarbonate formation. 2. Formation of bicarbonate. In this step, the amine behaves simply as a base, reacting with the carbonic acid formed from dissolution of CO2 in water. This step proceeds with a CO2: amine ratio of 1 and is associated with a relatively low enthalpy of absorption (and thus desorption). Obviously, the bicarbonate also exists in equilibrium (with carbonate and carbonic acid). The position of this equilibrium is dictated by pH and will affect the CO2:amine ratio. Under absorber conditions, however, the bicarbonate predominates. This pathway will be referred to in future as the acid pathway. 3. Formation of carbamate. Carbamate formation results in a CO2:amine ratio of 0.5. That is, two molecules of amine are required to absorb one molecule of CO2: one amine molecule is required for carbamate formation, while a second acts as a base to react with the proton released. This low CO2:amine ratio results in a lower efficiency and lower CO2 capacity than that seen for the acid pathway. In addition, carbamate formation is associated with a high enthalpy of reaction. Reversing this reaction in the stripper consequently requires the addition of large quantities of energy, making the entire process relatively energy intensive. These three individual steps are associated with a series of reactions as shown below: The concentrations of each species formed can be calculated based on the quantities of each initial component (amine and CO2) and the equilibrium constants of each chemical reaction involved. 1. Dissolution of the Gaseous CO2. Gaseous CO2 dissolves in water according to eq 1.

CO2(g) h CO2(aq)

(1)

Aqueous CO2 undergoes hydration to form carbonic acid, as shown in eq 2.

CO2(aq) h H2CO3

A + CO2(aq)* h ACOO- + H+

Neutralization Reactions. This group of reactions is not included in Scheme 1, as they do not directly involve CO2. However, these reactions are coupled to the reactions discussed above and cannot be omitted.

A + H+ h AH+

(7)

H+ + OH- h H2O

(8)

Temperature-dependent equilibrium constants (K) for the above equations have been reported previously in the form shown in eq 9.

ln K ) C1 +

C2 + C3 ln T + C4T T

(9)

where ln K is the natural logarithm of the equilibrium constant in terms of molality and T is the temperature in kelvin. The constants C1-C4 for each reaction are reported in the Supporting Information. This approximation provides a reliable representation of the temperature dependence of all equilibrium constants. The equilibrium constant for the protonation of MEA (KA) is available for a restricted temperature range only (0-50 °C, Bates et al.24), so it must be remembered that the extrapolation of this value is an assumption in the following model. All other constants are reported as valid to temperatures higher than those investigated in this work (temperature ranges are reported in the Supporting Information). Henry’s constant (eq 1) in 30 wt % aqueous MEA has been reported (Mandal et al.25), however only for temperatures between 20 and 40 °C, and does not differ significantly from the values reported in aqueous solution. Consequently, the values for Henry’s constant in aqueous solution (covering temperatures within the range 5-100 °C) were used. In order to calculate species concentrations, the activities of all species must be defined. In this work, the fugacity coefficients were fixed at unity as only atmospheric pressure conditions are considered in this study. The activity coefficients (γ) of each species in solution were calculated using an extended Debye-Hu¨ckel formula (eq 10) as used by Jou et al.26

(2)

The equilibrium for the hydration of carbonic acid lies significantly to the left and the relative concentration of H2CO3 is

(6)

log γi )

-Azi2xµ 1 + Baixµ

+2

∑j βijcj

(10)

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where ai is the ion size of species as given in Barner and Scheuerman;27 βij is the interaction parameter between the species i and j as determined by Jou et al.;26 cj is the concentration of species j; µ is the total ionic strength of the solution, given by µ ) (1/2)∑i cizi2, where ci is the concentration and zi is the charge of species i; A ) 50.3(T)-0.5, where T is the temperature in kelvin and  is the dielectric constant of the medium; and B ) (1.82 × 106)(T)-1.5. Temperature-dependent dielectric constants can be approximated according to eq 11.

 ) C5 + C6T + C7T2

(11)

The coefficients C5, C6, and C7 for water and MEA28 are given in the Supporting Information. The dielectric constant of a mixture is assumed to be additive. That is

(30% MEA) ) 0.3(MEA) + 0.7(H2O)

(12)

This is not strictly true as dissociation/protonation of the MEA and water will change their dielectric constants. The effect on the calculated solvent properties, however, will be minimal (see section entitled Solvent Properties under Modeling Novel Amine Systems), and this assumption is adequate for the work described herein. Given the above equilibrium constants and using activity coefficients as defined in eq 10, all equilibrium concentrations can be computed for any amine concentration, temperature, and any partial pressure of CO2 (up to 101 kPa, or 1 atm). In this case, the partial pressure of CO2 (PCO2) investigated is 12 kPa, as this is approximately the composition of flue gas from conventional coal-fired power plants. At this point, it must be noted that, in a real CO2 capture plant, the absorption column will not contain gas with a constant CO2 partial pressure of 12 kPa. At the bottom of the column, near the gas inlet, the partial pressure of CO2 will be close to 12 kPa. Higher up the column, some of the CO2 will have been removed from the gas stream and the partial pressure will be lower. Similarly, in the desorption column, the partial pressure of CO2 near the bottom of the column will be low. Toward the top of the column, the gas will contain more desorbed CO2, and the partial pressure will be higher. In this work, a simplified model is used, in which both absorber and desorber are assumed to operate at a CO2 partial pressure of 12 kPa. This model allows for easy comparison between systems. Although conditions are kept constant in this study, the model can easily be adapted to include concentration and temperature gradients throughout the absorber and stripper columns. This would give a more accurate picture of the chemistry inside a real capture facility. However, for the purposes of solvent comparisons discussed here, such a modification is unnecessary. From the equilibrium concentrations, the ratio (R) of absorbed CO2:amine can be calculated. Comparison of the ratios at absorption (40 °C) and desorption (100 °C) temperatures will give an indication of the capacity of the system. Enthalpies. The heat of reaction for each of the eqs 3-8 is assumed to be constant over the temperature range of interest, at the value determined at 25 °C. This value is calculated using the van’t Hoff equation (eq 13).

∆H )

∂ ln K 2 RT ∂T

(13)

The above assumption is not strictly true: heats of reaction are not totally temperature independent; however, over the tem-

perature range of interest, the deviation is typically less than 1%.29 Therefore, this is an adequate approximation for the work described herein. Given the equilibrium concentrations of each species, and the enthalpy associated with formation of that species, the enthalpy of absorption can be calculated. This is important for the validation of the model as the predicted values can be compared with previously experimentally determined ones. However, for the purpose of capture plant design and manufacture, the enthalpy associated with desorption, or the cyclic enthalpy, is of more relevance. In this work, the desorption enthalpy is calculated as the difference in absorption enthalpies under absorption and stripping conditions (40 and 100 °C, respectively), normalized to the number of moles of CO2 desorbed. Very few experimental determinations of desorption enthalpies have been undertaken. From the above discussion, it can be seen that, given a few basic equilibrium constants, the complete chemistry of CO2 capture by amine-based solvents can be modeled. The validity and accuracy of this model are demonstrated in the next section. Model Validation Equilibria. A. CO2 Solubility. The solubility of CO2 in aqueous amine systems is typically described in terms of the ratio of moles of CO2 taken up by the solution to the moles of amine in solution. The solubility of a 30 wt % MEA solution under different temperature and PCO2 conditions has been calculated using the model described above. The calculated CO2: MEA mole ratios (R’s) are in good agreement with those of Jou et al.26 and Lee et al.,30 as shown in Figure 2. Values are only calculated up to a CO2 partial pressure of 101 kPa as fugacity coefficients were fixed at 1, as discussed above. Higher pressures require different fugacity coefficients and were considered to be unnecessary as the capture plant typically operates at pressures close to atmospheric pressure and partial pressures of CO2 significantly lower than 101 kPa. The equilibrium constants used for the modeling computations are taken from a wide range of sources.23,24,26,28,31,32 The selected sources were those which covered a wide temperature range and were calculated based on activities, rather than concentrations. This explains, to some extent, the difference between the calculated and literature-reported results. One advantage of modeling the system in this way is that, in addition to total solubility, the speciation of the absorbed CO2 can be determined. That is, the relative quantities of absorbed CO2 present in different forms can be determined. As discussed in the Modeling section, the CO2 present as “acid” represents all CO2 present as part of the acid pathwaysdissolved CO2, carbonic acid, bicarbonate, and carbonate. The graphs in Figure 3 show the same calculated CO2 solubility values as shown in Figure 2 (s). It can be seen that at both temperatures, and all pressures, the CO2 is present predominantly as carbamate (- -). The proportion of CO2 present in acid form (---) only starts to become significant at high partial pressures of CO2. B. Cyclic Capacity. While the total absorption (or solubility) of the solvent in question is important, the cyclic capacity of the solvent is possibly more so. The cyclic capacity of the solvent is the difference between CO2 loadings in the rich and lean solutions. That is

cyclic capacity ) Rrich - Rlean

(14)

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Figure 2. Comparison of experimental solubility of CO2 in a 30% aqueous MEA solution with calculated results at 40 and 100 °C.

Figure 3. Distribution of absorbed CO2 at 40 and 100 °C. For clarity, the X and Y axes have been switched, compared to Figure 2.

Figure 4. Comparison of literature-reported and calculated enthalpies of CO2 absorption at 40 °C (left) and comparison of the desorption enthalpy (per mole of CO2 desorbed) with the absorption enthalpy (per mole of CO2 absorbed) (right). For better comparison of the two enthalpy values, the absorption enthalpy has been multiplied by -1.

where Rrich is the CO2:amine ratio in the CO2-rich solvent and Rlean is the CO2:amine ratio in the CO2-lean solvent. This value represents the capacity of the solvent over multiple cycles. In this study, R40 (CO2:amine ratio in the solvent at 40 °C) and R100 (CO2:amine ratio in the solvent at 100 °C) are substituted for Rrich and Rlean, respectively, to give an estimate for the cyclic capacity of the system. Assuming flue gas conditions of PCO2 ) 12 kPa and 30 wt % MEA, the cyclic capacity is 0.23. Once again, this can be split into a contribution from the carbamate pathway (0.20) and a contribution from the acid pathway (0.03). It is not surprising that the majority of the cyclic capacity is attributable to the carbamate, as it can be seen in Figure 3 that this is the major species formed. Enthalpy. At a PCO2 of 12 kPa, at 40 °C the enthalpy of initial CO2 absorption is -80 kJ/mol CO2 absorbed. Enthalpy of absorption is typically reported as a function of CO2 loading (ratio of CO2:amine), rather than as a function of CO2 partial pressure. The enthalpy of absorption calculated in this work is

shown as a function of loading in Figure 4a. The directly measured enthalpies of CO2 absorption33 are also shown (×), as are some values calculated from vapor-liquid equilibria (VLE) data (2, b).26,30 It is clear that the calculated enthalpy of initial absorption from this work is in good agreement with previously directly measured results, The enthalpy of absorption calculated in this way, however, differs significantly from the enthalpy calculated from VLE data, particularly at the higher ratios of CO2:MEA. This is unsurprising as enthalpies calculated from VLE data are typically only accurate to (30%.30 The enthalpy calculated above should not be confused with the enthalpy of repetitive absorption (or the cyclic enthalpy), as some CO2 will remain in solution after desorption, decreasing the amount of CO2 absorbed (and thus affecting the magnitude of the enthalpy) in subsequent cycles. This is most easily represented as the enthalpy of desorption, and will be constant for repeated cycles (assuming no degradation of the solvent). This value is of particular interest when designing capture plants

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Figure 5. Contributions of the four main reaction sets to the total enthalpy of CO2 desorption at 40 °C.

as it provides a realistic estimate of the energy requirements of the plant. The heat capacity of the solution is not incorporated in this work, but this value is easily incorporated into engineering models of the plant. In this work, the desorption enthalpy is described as a function of the moles of CO2 desorbed, and is calculated by

∆H100 - ∆H40 R40 - R100 ∆H100 - ∆H40 represents the absorption enthalpy difference between solutions at these two temperatures, and R40 - R100 equates to the number of moles of CO2 desorbed. For the MEA system investigated herein, the heat of desorption is approximately 83 kJ/mol CO2 desorbed. The enthalpy of absorption (per mole of CO2 absorbed) is compared to the enthalpy of desorption (per mole of CO2 desorbed) in Figure 4b. In the case of desorption, results are more clearly represented as a function of CO2 partial pressure. The difference between the two methods of reporting enthalpy becomes significant at higher partial pressures. As discussed above, modeling the system allows the speciation of the absorbed CO2 to be determined. This analysis can also be extended to the enthalpy of desorption. That is, the contributions of each individual reaction to the total enthalpy can be calculated, as shown in Figure 5. It is clear that a large contribution to the total enthalpy can be found in the neutralization step. The majority of this can, in turn, be attributed to protonation of the amine (eq 7). It is also worth noting that a significant contribution to the enthalpy of absorption is the dissolution of the CO2. This step must occur before either carbamate or bicarbonate formation can occur. The formation of the carbamate, however, contributes much more to the overall enthalpy than does the formation of bicarbonate/ carbonate. Given that the results are in good agreement with experimentally determined results, it can be assumed that this model adequately describes both the equilibrium conditions and the enthalpy of CO2 absorption. This model also returns enthalpies of absorption closer to directly measured results than those obtained indirectly through conventional calculations based on VLE data. Conditions in real absorbers, pilot-scale or large-scale CO2 capture plants, are different from those modeled above. As discussed above, a range of concentrations, pressures, and temperatures can be found along both the absorber and desorber columns. Typically, these columns are quantitatively described by compartmental models, in which the column is divided into sections with different conditions in each section. The capacity

and enthalpy properties of the individual sections are approximated with empirical expressions.21 The model described in this work can be used in place of the empirical approximations to more accurately calculate the properties of individual sections and thus build a complete model of a large-scale capture facility. Investigation into the Effect of Basicity and Carbamate Stability on Solvent Properties. Efforts are currently underway in several laboratories to develop improved solvent systems for CO2 capture.34,35 The model developed herein can be used to predict the effect of the amine properties (such as basicity and carbamate stability) on the performance of the solvent. In this way, the model can be used to guide the search for an alternative amine system. In this section, only simple monoamines are investigated. That is, they contain a single amine functionality and no other sites for protonation. The dielectric constant of the solvent was fixed at the value used for MEA, as was the Henry’s constant. Varying these parameters has little impact on the overall results, as discussed in the next section. In addition, the ionic radii of the amine and its carbamate were kept constant (again, the effect of this parameter is discussed in the next section). The parameters that were varied were the protonation constant of the amine (KA, associated with the reaction shown in eq 7) and the stability of the carbamate (Kcarb, associated with the reaction shown in eq 6). The plots in Figure 6 show the effect of changing these two parameters on both the enthalpy of desorption and the cyclic capacity of the system. The minimum desorption enthalpy occurs at the highest log KA and lowest log Kcarb values tested, while the maximum cyclic capacity occurs at approximately log KA ) 10.13 and the lowest log Kcarb value tested (cf. log KA ) 9.49 and log Kcarb ) -4.89 for MEA). Both graphs of Figure 6 illustrate that desorption enthalpy as well as cyclic capacity can be improved significantly with respect to MEA. The markers in the graphs indicate the position of MEA. A thorough analysis of the speciation changes at absorption and desorption temperatures would provide insight into the mechanisms behind the observed changes. This, however, is beyond the scope of this work and is not included herein. The increase in cyclic capacity with decreasing carbamate stability has been suggested previously36,37 and is discussed in the Introduction. The minimum calculated enthalpy within the range modeled is 31 kJ/mol CO2 desorbed, less than half the enthalpy of desorption of MEA. The maximum calculated cyclic capacity is 0.49, over double that of MEA. Clearly, in a simple monoamine solvent system, there is substantial room for improvement of the solvent, with respect to both enthalpy of reaction and cyclic capacity of the solvent. It must be remembered at this point that only a limited number of parameters were investigated herein. It is probable that further changes (such as incorporating additional amine groups onto the molecule, with associated acid dissociations and carbamate formations) will improve the properties of the absorption solution even further. Modeling Novel Amine Systems. This section addresses the challenge of modeling systems that have not been extensively investigated experimentally. This presents some challenges due to the fact that less information is available on new systems. Several parameters required for the model described above are unknown, or are estimates only. These issues are addressed below. A. Activity Coefficients. So far, the activity coefficients of all solution-phase species have been calculated using the extended Debye-Hu¨ckel formula as shown in eq 10. This

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Figure 6. Effect of changing the theoretical acid dissociation (KA) and carbamate stability (Kcarb) constants on the calculated desorption enthalpy (left) and cyclic capacity (right). The contour plots are shown below the surface plots for added clarity. The desorption enthalpy and cyclic capacity of MEA are marked with black circles on the contour plot. Table 1. Dielectric Constants of Some Simple Amines amine



temp (°C)

ref

triethylamine (NEt3) trimethylamine (NMe3) tributylamine (NBut3) diethanolamine (DEA) MEA

2.4 2.4 4.7 25.8 31.9

20 25 20 20 20

28 28 28 28 28

formula requires interaction parameters, βij. In the case of MEA, these interaction parameters were determined from a significant quantity of experimental data.26 In the absence of known interaction parameters, a further simplified activity coefficient formula must be used. Here, the formula used is the Davis equation (eq 15).

log γi ) - Azi2

(

xµ - 0.2µ 1 + xµ

)

(15)

In the case of MEA, comparing the results of applying the Debye-Hu¨ckel and Davis approximations, the difference in the enthalpy of absorption and the enthalpy of desorption is less than 1%, but the error in the cyclic capacity is almost 10%. The Davis equation, therefore, should only be used for comparative purposes and to provide an approximation of the cyclic capacity, rather than to give a precise value for this property. B. Physical Solubility of CO2. This model assumes that the physical solubility of CO2 in amine solutions is constant, at the same value as that of CO2 in water. The solubility of CO2 (Henry’s constant) in a variety of different amine solutions at different concentrations has been estimated.25 The minimum

reported value for Henry’s constant (at 40 °C) was -1.59 (2 M MEA), and the maximum reported value was -1.63 (3 M MDEA). These values are very close to those of water (-1.64), and this magnitude of change has a negligible effect on the calculated enthalpies and capacities associated with CO2 absorption. C. Solvent Properties. The dielectric constants () of both the amine and water are used to calculate the coefficients A and B in both the Davies (A only) and the extended DebyeHu¨ckel equations. The dielectric constant for MEA is significantly higher than that typically observed for amines.28 This can be attributed to the alcohol functionality and the low symmetry of the molecule. It is likely, therefore, that alternative alkanolamine systems will have similar dielectric constants. Several amines have been reported with dielectric constants below 5 (see Table 1), but dielectric constants above that of MEA are rare. To investigate the effect of changing this constant, the MEA system was modeled using the dielectric constant of triethylamine (NEt3). Differences (from the values obtained using the dielectric constant of MEA) are less than 1% for both calculated enthalpies and capacities. It is clear that the errors arising due to sensible estimation of the unknown parameters (or use of a slightly modified model) are small, and this model can be used to estimate both the heats of CO2 absorption and the capacities of an amine-based system. Further evidence for the suitability of this model for estimation of novel systems can be found in its applicability to alternative amine systems. Shown in Figure 7 are the previously reported and calculated enthalpy of absorption38,39 and solubility40,41 of CO2 at 40 °C in 50 wt % methyldiethylamine

Figure 7. Comparison of experimental solubility of CO2 in a 50% aqueous MDEA solution at 40 °C with calculated results (left) and comparison of literature-reported enthalpy in a 30% aqueous MDEA solution with calculated results (right).

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(MDEA). Values were calculated using the dielectric constant for MEA, the Henry’s constant for water, and the Davis equation (eq 15) to describe activity coefficients. In the case of MDEA, which contains a tertiary amine, carbamate formation does not occur. The temperature-dependent equilibrium constant for protonation of the amine (eq 7) was obtained from Kamps et al.32 (see Supporting Information for details). Excellent agreement with literature results can be seen for the solubility data. The agreement between the measured enthalpy and enthalpy calculated in this work is also reasonable, and the enthalpy calculated in this way is closer to the measured values than the enthalpy calculated from VLE data. Conclusions We have developed a model to describe the absorption of CO2 by aqueous MEA. This model has been used to accurately predict the enthalpy associated with both initial absorption of CO2 in aqueous MEA solution (80 kJ/mol CO2 absorbed) and desorption of CO2 from the resultant solution (83 kJ/mol CO2 desorbed). The model developed herein has also been used to accurately predict the CO2 solubility of the solution under different temperature and pressure conditions. The model was also used to predict a cyclic capacity of the solvent system of 0.23 mol of CO2/mol of amine, given absorber and stripper conditions of 40 °C and 100 °C, respectively, and a partial pressure of CO2 of 12 kPa. The CO2 speciation within the system was modeled, and it was shown that the majority of the CO2 captured by this system is present as carbamate. The effect of the amine basicity and the carbamate stability on both the cyclic capacity and the enthalpy of desorption have been investigated. It has been found that there is potential for substantial improvements in both the enthalpy of absorption and the cyclic capacity of the solvent system compared to MEA. Both properties can be improved by more than a factor of 2, compared to those of MEA. The effect of these properties is useful in directing the search for an amine system superior to MEA. The model has also been generalized for use in estimating the cyclic capacity and enthalpy of CO2 absorption for novel amine-based systems, for which only a minimal amount of information is available. We have demonstrated that the model can be used as a predictive tool for determining the performance both of novel amine systems and of different mixtures of existing solvent systems. This will allow rapid preliminary screening of new solvent systems based on faster and cheaper experiments (stability constant determinations) than the traditional vaporliquid equilibria experiments. Supporting Information Available: Coefficients for the temperature-dependent equilibrium equations as well as for the temperature-dependent dielectric constants. This material is available free of charge via the Internet at http://pubs.acs.org. Literature Cited (1) Stern, N. The Economics of Climate Change; www.hm-treasury.gov.uk: 2006. (2) Macfarling-Meure, C.; Etheridge, D.; Trudinger, C.; Steele, P.; Langenfelds, R.; van Ommen, T.; Smith, A.; Elkins, J. Law Dome CO2, CH4 and N2O Ice Core Records Extended to 2000 Years BP. Geophys. Res. Lett. 2006, 33, L14810. (3) Oreskes, N. Beyond the Ivory Tower: The Scientific Consensus on Climate Change. Science 2004, 206 (5702), 1686. (4) Pielke, R. A.; Oreskes, N. Consensus About Climate Change. Science 2005, 308, 952.

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ReceiVed for reView May 2, 2007 ReVised manuscript receiVed July 4, 2007 Accepted October 22, 2007 IE070619A