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Kinetics and Steric Hindrance Effects of Carbon Dioxide Absorption into Aqueous Potassium Alaninate Solutions Miri Kim, Ho-Jun Song, Min-Gu Lee, Ho-Yong Jo, and Jin-Won Park* Department of Chemical and Biomolecular Engineering, Yonsei University, 262 Seongsanno, Seodaemun-gu, Seoul 120-749, South Korea S Supporting Information *

ABSTRACT: Kinetic studies of carbon dioxide absorption into aqueous potassium alaninate (PA) solutions were performed using a stirred-cell reactor at concentrations ranging from 1.0 to 3.0 M and temperatures from 293.15 to 313.15 K. In the present work, the potassium salt of alanine was suggested as an energy-efficient absorbent for CO2 capture. Densities and viscosities of the aqueous PA solution and physical solubilities of CO2 in the solution were also measured to evaluate the reaction rate constant. The reaction rate was determined as the following equation: ⎛ − 3845 ⎞ ⎟exp(0.5706C )C C − rCO2 = 4.518 × 108exp⎜ s s CO2 ⎝ T ⎠

(−rCO2: rate of reaction between absorbent and CO2 gas [mol dm−3s−1]; T: temperature [K]; Cs: concentration of PA solution [mol dm−3]; CCO2: concentration of CO2 in PA solution [mol dm−3]). Furthermore, the effect of the sterically hindered structure of PA was studied by comparing the diffusivity, physical solubility, and reaction rate constant with literature values of potassium glycinate.

1. INTRODUCTION Global warming is threatening the survival of all human beings, causing climate change and ecocide. Of the greenhouse gases (GHGs), carbon dioxide (CO2) is the most urgent problem due to the enormous quantity (76% of total GHGs) being discharged by various industries including thermal power plants, iron, steel, and petrochemical industries.1 Under the postKyoto protocol regime, nations may need to spontaneously develop and commercialize CO2 capture and storage (CCS) technologies as a near-term solution to reduce CO2 emissions, because the present fossil fuel-dependent energy production and supply strategy cannot be altered in the near future. According to the International Energy Agency (IEA),2 CCS technologies could eliminate 19% of the global released CO2 until 2050. Among the various CO2 capture technologies, including adsorption, membrane separation, and cryogenics, the wet chemical absorption method using amine solution is thought to be the most feasible process to capture a large amount of CO2 that may be easily adopted by existing plants.3 Alkanolamines, such as monoethanolamine (MEA), diethanolamine (DEA), N-methyldiethanolamine (MDEA), and 2-amino-2-methyl-1-propanol (AMP), are extensively used as absorbents to selectively remove acidic gases (CO2, H2S) in gas streams. Although the alkanolamines have many advantages, they suffer from the following shortcomings that restrict commercialization of the wet CO2 absorption method: thermal and oxidative degradation, corrosiveness to the equipment, high absorbent regeneration energy, and high volatility.4−8 Recently, amino acid salts have been suggested as alternative CO 2 absorbents, because they offer advantages over alkanolamines. Amino acid salts have an identical functional group (−NH2) as © 2012 American Chemical Society

alkanolamines; hence, they can react with CO2 in a similar manner. Since amino acid salts have ionic structure in an aqueous solution, they offer lower volatility. Additionally, the salts have better resistance to oxidative degradation.9 A distinct characteristic of amino acid salt is the high surface tension of its aqueous solution, which makes it suitable as a gas absorbing liquid in a membrane-gas absorption (MGA) module. Some amino acid salts such as potassium salts of taurine and glycine are effective for MGA due to high membrane breakthrough pressure and fast reaction kinetics for CO2.10,11Additionally, amino acid salt solutions are more inexpensive than alkanolamine solutions, and production costs can be reduced considerably if amino acid salt solutions are used in a MGA process rather than a MEA-based process.12 Extensive studies of the reaction between CO2 and sodium glycinate (SG) amino acid salt, including reaction kinetics, thermodynamic properties measurements, and process simulation, were previously conducted by our group.13−18 Kumar et al.6,10,19−21 researched the aqueous potassium salt of taurine and reported its application to MGA, including important physical properties, CO2 solubility, and reaction kinetics. Additionally, they suggested a correlation between the solubility of amino acid in pure water and the critical solution concentration for which a precipitate occurs during CO2 absorption. Van Holst et al.22 selected promising CO2 absorbents, specifically proline and sarcosine among various amino acids (6-aminohexanoic acid, Received: Revised: Accepted: Published: 2570

July 26, 2011 December 8, 2011 January 10, 2012 January 10, 2012 dx.doi.org/10.1021/ie201609b | Ind. Eng.Chem. Res. 2012, 51, 2570−2577

Industrial & Engineering Chemistry Research

Article

β-alanine, L-arginine, L-glutamic acid, DL-methionine), by plotting the overall reaction kinetic constant (kov) against basicity (pKa) at a specific concentration and temperature. They also measured the apparent rate constants (kapp) over various concentrations of aqueous potassium and lithium salts of proline and sarcosine at 298 K. Portugal et al.23,24 conducted kinetic studies of CO2 absorption into aqueous potassium salts of glycine and threonine, which is the hydroxylated form of glycine. Simons et al.25 measured kinetic constants for the partially CO2-saturated aqueous solutions and CO2-lean solutions of potassium sarcosinate for more practical purposes. This research showed that certain amino acid salt solutions could be fine alternatives of alkanolmines. Kinetics of the reaction between CO2 and aqueous solutions of α-alanine were conducted in the present study. As shown in Figure 1a, α-alanine is a sterically hindered amino acid, in which

CO2 reacts with potassium alaninate and produces a zwitterion. Then, some of the zwitterions are deprotonated by bases (Bi), such as amine, water, and OH−. Assuming a quasi-steady-state condition for the zwitterion concentration, the overall CO2 absorption rate can be given by

− rCO2 =

k2 1+

k −1 ∑i k BiC Bi

CsCCO2 (3)

where −rCO2 is the rate of reaction between amino acid salt and CO2 [mol dm−3 s−1], CS is the concentration of the amino acid salt solution [mol dm−3], and CCO2 is the concentration of CO2 in the solution [mol dm−3]. The effect of OH− on the deprotonation of the zwitterion is negligible,29 and the rate of zwitterion deprotonation is faster than the reverse reaction of 1 (k−1/∑ikBiCBi ≪ 1).23 Then, eq 3 is reduced to

− rCO2 = k 2CsCCO2

In principle, the CO2 absorption rate must be expressed by activities instead of concentrations for thermodynamic consistency,30 but it requires equilibrium data that are not available. To account for the nonidealities of the solution, the following semiempirical equation is generally applied,31

Figure 1. Chemical structures of (a) alanine and (b) potassium salt of alanine (PA).

the electron donating-methyl group is attached to α-carbon. Thus, easy absorbent regeneration and high CO2 absorption capacity are expected.9 Furthermore, the methyl group between the amino (−NH2) and carboxylic (−COOH) groups effectively blocks the attack of V5+, which is from the corrosion inhibitor (V2O5), so that the degradation of the amino group would be prevented.26 In our preliminary experiment, the surface tension of the aqueous potassium alaninate (CAS number: 93893-38-0) solution increased as the concentration increased, indicating that potassium alaninate (PA) solution could be utilized in MGA without causing the membrane porewetting problem. The aqueous PA solutions were prepared by adding equimolar KOH (Figure 1b) to alanine in order to recover the reactivity of the amino group in water. Kinetic studies were performed using a stirred-cell reactor (SCR), and the steric hindrance effect was briefly discussed via comparisons of the CO2 absorption rates between PA and potassium glycinate (PG). The data presented in this study will be used as engineering data for the design of a PA-based acid gas removal (AGR) unit in the future.

keff = k exp(bI )

NCO2 = E ·kL

E=− +

(1)

A (6)

Ha2 2(E∞ − 1) Ha 4 4(E∞ − 1)2

+

E∞·Ha2 +1 E∞ − 1

(7)

The Hatta number is defined as

−

OOC+H2N − CH(CH3) − COO−K+ + Bi ⎯→ ⎯ −OOCHN − CH(CH3) − COO−K+ + Bi H+

HCO2

where kL is the physical mass transfer coefficient [m s ], PCO2 is the CO2 partial pressure in the gaseous phase [mbar], HCO2 is the Henry constant of CO2 in the solution [kPa m3 mol−1], A is the interfacial area between the gas and liquid phases [m2], and E is the enhancement factor. The enhancement factor, the ratio between the absorption rates, increases by the chemical reaction, and the rate of physical absorption is a function of the Hatta number (Ha) and the infinite enhancement factor (E∞) [dimensionless] as in the eq 7.33

k2

k Bi

PCO2

−1

CO2 + H2N − CH(CH3) − COO−K+ k −1

(5)

where keff is the effective kinetic constant [dm3 mol−1 s−1], which is reflecting the ionic strength of the solution, b is a constant, and I is the ionic strength [mol dm−1] given by I = (1/2)∑iCizi2 where Ci is the molar concentration [mol dm−3] and zi is the charge number of ions in the solution. For low CO2 loading and both monovalent species (z2 = 1), I becomes similar to Cs. 2.2. Mass Transfer. The molar flow of carbon dioxide into lean potassium alaninate solution is described by the following equation,32

2. THEORY 2.1. Zwitterion Mechanism. The zwitterion mechanism was originally proposed by Caplow27 and later modified by Danckwerts.28 According to this mechanism, the following equations can be established for potassium alaninate salt:

←→ −OOC+H2N − CH(CH3) − COO−K+

(4)

Ha = (2) 2571

kovDCO2 kL

(8)

dx.doi.org/10.1021/ie201609b | Ind. Eng.Chem. Res. 2012, 51, 2570−2577


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where kov is the overall reaction kinetic constant [s−1] and DCO2 is the diffusion coefficient of CO2 in the solution [m2 s−1]. According to the penetration model,34 the enhancement factor, E∞, is given by

E∞ =

DCO2 DS

+

CS νCO2

PCO2 HCO2

DS DCO2

mass fraction was accurate to ±0.5%. H2SO4 was purchased from Duksan Chemical Co. 3.2. Density and Viscosity. The densities of the aqueous potassium alante solutions were measured at 293.15, 298.15, 303.15, and 313.15 K using a calibrated density meter (DMA 35, Anton Paar). The viscosities were also measured using a commercial viscometer (RheoStress1, HAAKE Instruments Inc.) with a DG 41 spindle at identical temperatures. 3.3. Physical Solubility. Physical solubility cannot be directly measured, as amino acid salt solutions react with carbon dioxide. In this study, N2O was used to measure these properties indirectly because N2O has similar molecular structure, electric structure, and volume to that of CO2, and it is not reactive with the amino acid salt.35 The experimental apparatus (stirred cell reactor) is similar to that of Kumar et al.21 and was composed of two vessels with calibrated volumes; one for storing the N2O gas and the other for the absorbent solution with two impellers inside for the gas and liquid phases. Then, 470 mL of the PA solution was fed into the 980 mL reactor, which was evacuated forcibly with the vacuum pump “until the pressure reached about −0.970 bar” (275 mini Convectron, Helix Technology Corp.). The vapor equilibrium was allowed to reach a given temperature (293.15, 298.15, 303.15, and 313.15 K), and the vapor pressure (Pvapor) was then recorded. The gas vessel was filled with N2O. A known amount of N2O (about 750 mbar) was allowed to enter the absorbent vessel, and the initial pressure (Pinit) was recorded. The stirrer was switched on, and the solution equilibrium was allowed to be established. The final equilibrium pressure (Peq) and temperature (Tinit) were recorded. The temperature was set to a different value, T, by thermostatic bath (RBC-11, LAB House), and a new equilibrium state was established. The amount of absorbed gas can be calculated by applying the ideal gas law. Additionally, the Henry coefficient for N2O (HN2O) is then computed with the following equation:23,24

(9)

where Ds is the diffusion coefficient of the amino acid salt and νCO2 is the stoichiometric coefficient. The infinite enhancement factor changes as the partial pressure of carbon dioxide changes inside the reactor, and consequently, the ratio between Ha and E∞ changes, which means that the absorption regime changes. When the ratio between the Ha and E∞ is relatively large (3 < Ha ≪ E∞), the reaction of CO2 with the amino acid salt occurs in the pseudofirst-order (PFO) regime32 and the enhancement factor equals Ha. Hence, eq 6 becomes:

NCO2 =

kovDCO2

PCO2 HCO2

A (10)

Additionally, if the PFO regime is fulfilled, k2Cs in eq 4 can be defined as kapp, which is the apparent kinetic constant between the absorbent and CO2. Because the contribution of the reaction between CO2 and hydroxide ions is negligible, the kapp becomes comparable to kov, which is the overall kinetic constant (kapp = kov + kOH−[OH−] ≈ kov).29

3. EXPERIMENTAL SECTION 3.1. Chemicals. Aqueous potassium alaninate solutions were prepared by adding an equimolar quantity of potassium hydroxide (Acros, 8 N aqueous solution) to alanine (Alfa Aesar, ≥ 99%), which was dissolved in demineralized water. The mass fraction of potassium alaninate in the aqueous solution was determined by titrating against 0.1 M H2SO4. The measured HN2O(T ) =

⎛ VL ⎞ ⎜ ⎟ Pinit − Pvapor(Tinit) Peq(T ) − Pvapor(T ) ⎝ RV ⎠ G − Peq(T ) − Pvapor(T )

Tinit

T

where VG and VL are the volume of gas and liquid in the absorbent vessel [m3], respectively, and R is the universal gas constant, 8.314 J mol−1 K−1. The solution vapor pressure at each temperature was estimated by the following Raoult’s relation:

Pvapor(T ) = x H2OPHpure (T ) 2O

(11)

reactor records the internal pressure drop due to the CO2 absorption reaction. If the PFO regime is fulfilled, the overall reaction constant is determined by the following equation:36 ln(PCO2,t) = ln(Pt − Pvapor) = ln(P0 − Pvapor) −

(12)

kovDCO2

RT A·t VGHCO2

(13)

where VG is the volume of the gas phase in the reactor [m3] and A is the contact area between the liquid and gas phase [m2]. Although the experimental process in the batch system is simple and clear, it is somewhat inaccurate compared to the semibatch analysis, especially for absorbents with fast absorption rates. This is because CO2 is consumed rapidly before the system reaches the PFO regime. However, in this study, the potassium alaninate solution had a constant slope in a specific pressure range because of its slow absorption rate. In order to enhance experimental accuracy, several experiments were performed for each temperature and concentration and analyzed using statistical software. Each experiment was processed during the CO2 pressure drop from approximately

All of the initial pressures and a data sheet are provided in the Supporting Information. 3.4. CO2 Absorption Kinetic Measurements. Similar to the N2O solubility experiment, 470 mL of PA solution was fed to the absorbent vessel of the stirred cell reactor and trace gases were removed by applying a vacuum. The vapor equilibrium was allowed to reach a given temperature, and the vapor pressure (Pvapor) was recorded. Instead of N2O gas, CO2 was allowed to enter the absorption vessel for the measurement of CO2 absorption flux into the PA solution. After a certain amount of CO2 gas is injected, the valve was closed to make the reactor a batch system. A computer program (Cimon X, KDT Systems Co.) connected to the pressure indicator of the stirred 2572



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100 mbar to less than 5 mbar (20−30 min). Figure 2 is a representative sample of this experiment (PA solution of 2 M

Table 2. Sechenov Constants and Specific Parameters of the Schumpe Model for N2O and CO2 in a 1.0 M PA Solution

293.15 298.15 303.15 313.15

0.456698641 0.485834304 0.5176258 0.613444486

−0.0061 −0.0085 −0.0109 −0.0157

−0.0155 −0.0172 −0.0189 −0.0223

0.437898641 0.468434304 0.5016258 0.600244486

D N2O D N2O,w

(17)

=

DCO2 DCO2,w

(18)

Diffusivities of N2O and CO2 ranging from 1.0 to 3.0 M are presented in Figure 4. 4.4. Kinetic Measurements. Initially, the ratio of the flux caused by the chemical reaction at the liquid−gas interface was greater than the flux caused by physical diffusion because the reaction occurs only at the liquid−gas surface and it takes some time for the reactant to diffuse into the reaction zone in the liquid. In this instantaneous regime, the enhancement factor E is presented as E∞ and E∞ ≪ Ha. After the chemical reaction is saturated (reaches equilibrium) at the interface, CO2 gas physically penetrates into the liquid phase. Therefore, as the reaction proceeds, the fraction of flux caused by physical diffusion among the total absorption flux increases and the chemical reaction between CO2 and the absorbents becomes the rate-determining step (E = Ha, 3 < Ha ≪E∞, PFO regime). Although the stirrer in the reactor rotated continuously, a uniform solution was not possible because the rotation rate was fixed below 60 rpm in order to maintain the liquid−gas interface area where CO2 absorption occurs constant. kov, determined by the slope of the PFO regime of the plot and eq 13, is described in Table 4. Applying the low CO2 loading condition, both monovalent species (z2 = 1) in eq 5, and the

(14)

where HN2O and HN2O,w are the Henry’s constants of N2O in the amino acid salt solution and water [kPa m3 mol−1], respectively, and K is the Sechenov constant [dm3 mol−1]. For a single salt, K can be calculated theoretically by the following equation:38 (15)

where hi and hG are the ion and gas specific parameters [dm3 mol−1], which are given in Table 2, and ni is the ion’s valence number, where nAla− = 1 and nK+ = 1. Then, Sechenov constant becomes24

K CO2 = K N2O − 2h N2O + 2hCO2

KCO2 (dm−3 mol)

where α is a constant that depends on the pair gas/solution. For an amino acid salt solution, α = 0.6.39 The diffusion coefficient of CO2 in solution is then determined using the N2O/CO2 analogy:41

4. RESULTS AND DISCUSSION 4.1. Density and Viscosity. The densities and viscosities of aqueous potassium alaninate solutions at temperatures from 273.15 to 313.15 K and concentrations from 1.0 to 3.0 M are presented in Table 1. 4.2. CO2 Physical Solubility. Henry’s constant is influenced by solution concentration due to the effect of ion salting out of amino acid salt. To consider this effect, the solubility data of N2O in potassium alaninate was fitted using the following Sechenov relation, which was developed by Schumpe,37

∑ (h i + h G )n i

hCO236 (dm−3 mol)

D·ηa = constant

at 298.15 K), and a data sheet is shown in the Supporting Information.

K=

hN2O36 (dm−3 mol)

By knowing KCO2, the physical Henry’s constant of CO2 (Hp,CO2) can be obtained using the Sechenov relation for CO2. The parameters and Sechenov constants are suggested in Table 2, and physical Henry’s constants are presented in Table 3 and Figure 3. 4.3. Gas and Ion Diffusion Coefficients. It is difficult to directly determine the diffusion coefficient because it is very time-consuming and requires complicated calculation formulas. A modified Stokes−Einstein equation is generally applied to indirectly estimate diffusivity.19,39,40 The modified Stokes−Einstein equation is as follows:

Figure 2. Representative sample of CO2 absorption kinetic experiment (for 2 M PA solution at 298.15 K).

⎛ HN O ⎞ 2 ⎟⎟ = K · Cs log⎜⎜ H ⎝ N2O,w ⎠

T (K)

KN2O (dm−3 mol)

(16)

Table 1. Densities and Viscosities of Aqueous PA Solutions T (K)

293.15 −3

298.15 −1 −1

−3

303.15

ρ (kg m )

η × 10 (kg m s )

ρ (kg m )

η × 10 (kg m s )

ρ (kg m )

η × 10 (kg m s )

ρ (kg m )

η × 103 (kg m−1 s−1)

1.0 2.0 3.0

1055 1109 1160

1.507 2.173 3.291

1053 1107 1158

1.359 1.928 2.927

1052 1105 1156

1.224 1.719 2.590

1048 1100 1151

1.041 1.412 2.082

3

−1 −1

2573

−3

313.15

Cs (M)

3

3

−1 −1

−3



Article

Table 3. Physical Henry’s Constants of N2O and CO2 in Water and a PA Solution Hp,CO2 (Pa·m3·mol−1) T (K)

293.15

298.15

303.15

313.15

Cs (M)

N2O

CO2

N2O

CO2

N2O

CO2

N2 O

CO2

water 1.0 2.0 3.0

3492 4913 6647 9188

263237 3546 4594 6082

4069 5716 7370 10102

304037 4103 5082 6692

4605 6529 8143 11148

357137 4881 5867 7741

5720 8129 9922 13129

421937 5816 6887 8840

Table 4. Overall Kinetic Constants of the Reaction between the PA Solution and CO2 kov (s−1) Cs (M)

293.15 K

298.15 K

303.15 K

313.15 K

1.0 2.0 3.0

2292 5222 15303

3109 6135 18157

3777 7652 23982

5202 12577 34587

Figure 3. CO2 physical Henry’s constants of 1.0−3.0 M PA solutions at temperatures from 293.15 to 313.15 K.

Figure 5. Experimental and calculated values of kov.

The reason for the exceptionally large error between the experimental and calculated data for the 1 M PA solution is thought to be the correction for nonideality in eq 19. As described previously, eq 19 is derived through considerations for the nonideality of a solution (see eq 5). For this reason, further error may be encountered for a low concentration solution, which is similar to an ideal solution. Also, according to Portugal et al., 24 experimental data are somewhat different from modeled values and the difference is intensified for low concentration potassium threonate solutions (under 1 M). From eqs 4 and 19, the final equation is derived as follows:

Figure 4. CO2 diffusivity into 1.0−3.0 M PA solutions at temperatures from 293.15 to 313.15 K.

Arrhenius equation, the following equation is obtained.

⎛ − 3845 ⎞ ⎟exp(0.5706C )C kov = 4.518 × 108exp⎜ s s ⎝ T ⎠

⎛ − 3845 ⎞ ⎟exp(0.5706C )C C − rCO2 = 4.518 × 108exp⎜ s s CO2 ⎝ T ⎠

(19)

(20)

where −rCO2 is the rate of reaction between absorbent and CO2 gas [mol dm−3s−1] and CCO2 is the concentration of CO2 in PA solution [mol dm−3]. 4.5. Steric Hindrance Effect. PA is a steric hindered form of PG, having a methyl group at its α-carbon. Experimental and theoretical data of PA and literature PG data are compared and briefly discussed. 4.5.1. CO 2 Diffusivity and Physical Solubility. CO 2 diffusivity and physical Henry’s constants ranging from 1.0 to

(kov: overall kinetic constant [s−1]; T: temperature [K]; Cs: concentration of PA solution [mol dm−3]). Comparisons between the experimental values in Table 4 and the calculated line by eq 19 are shown in Figure 5. As the temperature and concentration of the PA salt increased, the overall kinetic constants increased exponentially. The error bars for the experimental values seem to indicate acceptable tolerance, even considering the fact that the y-axis of this figure is on a logarithmic scale. 2574



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3.0 M are presented in Tables 5 and 6. CO2 diffusivity of PA is less than that of PG by approximately 10% according to Table 5, which means that PA is more resistant to the CO2 Table 5. Comparison of CO2 Diffusivity into PA and PG21 Solutions DCO2 × 109 (m2·s−1) T (K)

293.15

298.15

303.15

313.15

Cs (M)

PA

PG21

PA

PG21

PA

PG21

PA

PG21

1.0 2.0 3.0

1.32 1.06 0.83

1.42 1.16 0.94

1.51 1.23 0.96

1.58 1.30 1.07

1.83 1.50 1.17

1.80 1.48 1.20

2.14 1.79 1.42

2.24 1.82 1.50

diffusion. Additionally, the Henry’s constant of CO2 is greater than the PG solution at PA by approximately 6−7%, indicating that the physical solubility of PA is less than PG. These experimental results may be explained by comparing the molecular structures of these two materials. Compared to PG, PA has a substituted bulky methyl group on its α-carbon. The larger volume caused by this methyl group may increase the molecular interactions between PA salts by the addition of van der Waals forces, consequently causing the increase of the solution viscosity. CO2 diffusivity into the PA solution is then less than that into the PG solution based on eq 17. This interpretation is consistent with the increased Henry’s constant, meaning decreased physical CO2 solubility. 4.5.2. CO2 Flux. A comparison of the PA and PG solution kinetics is shown in Figure 6. The kov of the PA solution is less than that of the PG solution. Competitive aspects of the steric hindrance effect on the reaction between the absorbent and CO2 occur. The first is that the bulk groups of the sterically hindered absorbents, the methyl group at the α-carbon, physically obstruct CO2 from access to the N atom, and consequently, the reaction rate becomes slow. The other aspect is that nucleophilicity of N atom is increased by the tendency of the methyl group and N atoms to give electron pairs to the C atom of CO2, thereby increasing its chemical reactivity.42 According to Jhon et al.,43 the charge density of the N atom of the AMP carbamate ion (0.128 e) is less than that of the MEA carbamate ion (0.144 e), which is the nonsterically hindered form of AMP. As a result, the AMP forms a more stable C−N bond by electron delocalization and its CO2 absorption capacity considerably increases. However, in terms of the collision frequency, MEA has a higher value than AMP because of the physical disturbance of the AMP bulk group. In this study, the CO2 absorption reaction kinetic of the PA solution is much slower than that of the PG solution, suggesting that the effect of the increase of the nucleophilicity by the steric hindered structure is not more critical than the physical hindrance effect.

Figure 6. Comparison of kov for PA and PG solutions.

5. CONCLUSIONS We evaluated important physical properties (density, viscosity, Henry’s constant and diffusivity of CO2 gas) and CO2 absorption kinetics under the PFO regime of a PA solution. Experiments were performed in a stirred cell reactor at concentrations ranging from 1.0 to 3.0 M and temperatures from 293.15 to 313.15 K. These results are compared with PG solution data from the literature for consideration of the steric hindrance effect. The PA solution has a lower CO2 absorption rate than the PG solution. However, the steric hindrance effect of the PA solution might improve its CO2 loading capacity by increasing the charge density of the N atom of the PA salt. CO2 absorption kinetics were closely associated with the desorption rate, and these two factors show reciprocal behavior.44 In the case of the PA solution, the fast desorption rate caused by slow absorption rates decreased the amount of energy consumed, with recovery of CO2 from the absorbents to improve its regeneration efficiency.

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ASSOCIATED CONTENT

S Supporting Information *

Tables of N2O solubility and CO2 absorption kinetic experiment. This material is available free of charge via the Internet at http://pubs.acs.org.

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AUTHOR INFORMATION

Corresponding Author

*Tel.: +82-2-364-1807. Fax: +82-2-312-6401. E-mail: jwpark@ yonsei.ac.kr.

Table 6. Comparison of Henry’s Constant of CO2 into PA and PG21 Solutions Hp,CO2 (Pa·m3·mol−1) T (K)

293.15

Cs (M)

PA

1.0 2.0 3.0

3546 4594 6082

298.15 PG

21

3340 4212 5313

PA 4103 5082 6692

303.15 PG

21

3725 4662 5835 2575

PA 4881 5867 7741

313.15 21

PG

4138 5139 6382

PA

PG21

5816 6887 8840

5053 6179 7554



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ACKNOWLEDGMENTS This work was supported by the National Research Foundation of Korea (NRF) grant funded by the Korea government (MEST) (No. 2011-0029161).

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NOMENCLATURE A = gas−liquid interfacial area, m2 C = concentration, M (mol dm−3) D = diffusivity coefficient, m2 s−1 E = enhancement factor, dimensionless E∞ = infinite enhancement factor, dimensionless h = ion and gas specific constants in the Schumpe equation, dm3 mol−1 Hp = physical Henry’s constant, kPa m3 mol−1 Ha = Hatta number, dimensionless I = ionic strength of the solution, mol dm−1 k−1 = kinetic constant for reverse reaction of zwitterion formation, s−1 k2 = reaction kinetic constant for zwitterion formation, dm3 mol−1 s−1 kapp = apparent kinetic constant, dm3 mol−1 s−1 kBi = kinetic constant for zwitterion deprotonation by base, dm3 mol−1 s−1 kOH− = kinetic constant for the reaction between CO2 and OH−, dm3 mol−1 s−1 kov = overall kinetic constant, s−1 K = Sechenov constant, dm3 mol−1 NCO2 = CO2 absorption flow, mol s−1 PCO2 = partial pressure of CO2, mbar −rCO2 = rate of reaction between amino acid salt and CO2, mol dm−3 s−1 R = universal gas constant, 8.314 J mol−1 K−1 T = temperature, K V = volume, m3 α = constant in the modified Stokes−Einstein equation η = viscosity of solution, kg m−1 s−1 υCO2 = stoichiometric coefficient ρ = density of solution, kg m−3 z = charge number, dimensionless

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