Ab Initio Study of CO2 Capture Mechanisms in Aqueous 2-Amino-2

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Ab Initio Study of CO2 Capture Mechanisms in Aqueous 2Amino-2-methyl-1-propanol: Electronic and Steric Effects of Methyl Substituents on the Stability of Carbamate Yoichi Matsuzaki, Hidetaka Yamada, Firoz Alam Chowdhury, Shin Yamamoto, and K. Goto Ind. Eng. Chem. Res., Just Accepted Manuscript • DOI: 10.1021/acs.iecr.8b06229 • Publication Date (Web): 07 Feb 2019 Downloaded from http://pubs.acs.org on February 15, 2019

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Industrial & Engineering Chemistry Research

Ab Initio Study of CO2 Capture Mechanisms in Aqueous 2-Amino-2-methyl-1-propanol: Electronic and Steric Effects of Methyl Substituents on the Stability of Carbamate Yoichi Matsuzaki,† Hidetaka Yamada,‡ Firoz A. Chowdhury,‡ Shin Yamamoto,‡ and Kazuya Goto‡

†Advanced

Technology Research Laboratories, Nippon Steel & Sumitomo Metal Corporation, 20-1

Shintomi, Futtsu, Chiba 293-8511, Japan ‡Chemical

Research Group, Research Institute of Innovative Technology for the Earth, 9-2

Kizugawadai, Kizugawa, Kyoto 619-0292, Japan

To whom correspondence should be addressed. E-mail: [email protected].

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Abstract

Ab initio quantum chemical calculations, combined with the SMD solvation model, have been conducted to reveal CO2 absorption mechanisms and the origin of the reduced stability of carbamate in aqueous solution of 2-amino-2-methyl-1-propanol (AMP). The calculated Gibbs free energy change for all the elementary steps in the conversion of carbamate to bicarbonate reveal that the reduced thermodynamic stability of AMP carbamate is largely ascribable to the shift of the equilibrium from the carbamate to the zwitterion intermediate, from which CO2 can be released and successively converted to bicarbonate. This feature of AMP is an attribute of the electronic effects of the substitution by methyl groups specifically at the -carbon. Furthermore, the steric effects of the methyl groups are investigated from a dynamical point of view that suggests a new scenario in which the methyl groups promote the release of CO2 from the zwitterion intermediate.


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INTRODUCTION There is a consensus that the emission of greenhouse gases, including CO2, is the main cause of global warming and climate change problems.1 Chemical absorption using an aqueous amine solution is currently one of the most promising technologies to capture CO2 from large point exhaust gas streams.2 In practical applications of amine-based CO2 capture technology, it is necessary to reduce the capture cost, the majority of which arises from absorbent regeneration, i.e. the release of CO2 from the aminesolvent solution. Therefore, it is essential to reduce the energy consumption due to absorbent regeneration by developing new absorbents with preferred features like high absorption rate, high absorption capacity, and low heat of reaction. To this end, it is crucial to understand the complete reaction mechanism underlying the CO2 capture process, based on the molecular structures of the amines. In aqueous solutions of primary or secondary amines, CO2 is absorbed by the formation of carbamate: (1)

CO2 + 2RNH2 ⇄ RNHCOO + RNH3+.

In the above reaction, we have considered only primary amines, because we focus only on them in this study. The reaction in Eq. 1 is now widely accepted to proceed through a zwitterion intermediate, as represented by the combination of the reactions in Eq. 1a and 1b.3-5 CO2 + RNH2 ⇄ RNH2+COO,

(1a)

RNH2COO + RNH2 ⇄ RNHCOO + RNH3+.

(1b)

CO2 can be also absorbed in the form of bicarbonate: CO2 + H2O + RNH2 ⇄ HCO3 + RNH3+.

(2)

The net effect of the reactions in Eq. 1 and 2, i.e., (2) - (1), formally represents the hydrolysis of carbamate: (3)

RNHCOO + H2O ⇄ HCO3 + RNH2. It is to be noted that the direct reaction in Eq. 3 is kinetically hampered

6,7

while it represents an


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overall equilibrium between carbamate and bicarbonate. Generally, the reactions in Eq. 1 and 2 can proceed with primary or secondary amines in aqueous solutions. However, since the molar ratio of CO2 and amine is 1:2 in Eq. 1 and 1:1 in Eq. 2, CO2 absorption in the form of bicarbonate is preferred to that in the form of carbamate to increase the CO2 loading capacity. In addition, amines that prefer the formation of bicarbonate to carbamate tend to have lower heats of reaction. In this regard, tertiary amines absorb CO2 exclusively in the form of bicarbonate, but suffer from low absorption rates. An alternative way to increase the amount of bicarbonate, without significantly decreasing the reaction rate, is to introduce bulky substituents close to the nitrogen atom in primary or secondary amines. This results in the so-called sterically hindered amines.8 2-Amino-2-methyl-1-propanol (AMP, Figure 1) is a representative sterically hindered amine.8 NMR measurements have revealed that an aqueous solution of an unhindered amine, monoethanolamine (MEA, Figure 1), absorbs CO2 predominantly in the form of carbamate, while an aqueous solution of AMP predominantly undergoes bicarbonate formation.8,9 NMR speciation10,11 and other12 studies have been performed to determine the carbamate stability constant Kc and is defined as (4)

Kc ≡ [RNHCOO]/[RNH2][HCO3]

According to 1H NMR speciation, the MEA to AMP ratio of Kc is larger than 174 at 30 °C, thereby demonstrating a remarkable difference in the form of CO2 captured at equilibrium.10

H2N

OH

H2N

MEA

OH

AP H2N

H2N

PA

H2N

OH

AMP

H2N

SBA

TBA

H2N

OH

AMPb H2N

IBA

Figure 1. Amines investigated in this study. ACS Paragon Plus Environment

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In order to explain the characteristic CO2 capture by aqueous AMP solution, it has been considered that AMP carbamate is unstable because of the steric hindrance from the methyl groups8 and can be readily hydrolyzed, as represented by the reaction in Eq. 3. However, earlier theoretical studies have attributed the reduced stability of AMP carbamate to the electronic effects of the methyl groups at the carbon atom– the amino nitrogen atom becomes a softer base, thus reducing its affinity toward the hard acid, CO2.13 Indeed, the C-N bond length in the X-ray crystal structure of AMP carbamate does not exhibit significant elongation,14 indicating that the steric repulsion from the methyl groups is not so significant. Furthermore, our quantum chemical calculations6,7 have revealed that the energy barrier for the reaction in Eq. 3 is considerably high for both MEA and AMP, thus eliminating the possibility of Eq. 3 being an elementary reaction. Because of these findings, a further investigation is required to reveal the underlying mechanism for the CO2 capture by sterically hindered amines. Recent ab initio molecular dynamics simulations15 have demonstrated the influence of solvation around the amino nitrogen on the kinetics of the reactions in Eq. 1 and 2. It has been predicted that the strong solvation around the nitrogen atom of AMP prohibits the access of CO2 and promotes the deprotonation of H2O to OH, which attacks CO2 to produce bicarbonate. The influence of micro solvation on the kinetics has also been proposed in ref 16. In this study, we performed ab initio quantum chemical calculations to reveal the underlying mechanisms for the preferred bicarbonate production from CO2 capture in aqueous AMP solution. We reveal the electronic effects of the methyl groups on the thermodynamic and kinetic stabilities of AMP carbamate. The unique feature of AMP is highlighted by comparing its behavior with other amine compounds whose geometrical structures have been systematically varied (Figure 1).10 Furthermore, the steric effects of the methyl groups are investigated from a dynamical point of view that suggests a new scenario in which the methyl groups promote the release of CO2 from the zwitterion intermediate.

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All the calculations were carried out using the Gaussian09 suite of program.17 The molecular geometries were optimized using the Møller-Plesset second-order perturbation (MP2) theory combined with the 6-311++G(d,p) basis set, and all the optimized geometries were confirmed to be located on the minimum points on the potential energy surfaces by performing normal vibration analyses. On the basis of the optimized geometries, single-point energy calculations were performed using the coupled-cluster theory with single, double, and non-iterative triple excitation (CCSD(T)), combined with the 6311++G(2df,2p) basis set. In all the calculations, solvent polarization effects were incorporated by using the SMD18 variant of the continuum solvation model (CSM) with the dielectric constant of water (78.39). It is well known that the CSM has a problem in describing explicit solute-solvent interaction such as hydrogen bonds which plays a significant role in stabilizing ionic species. Including explicit solvent molecules around the solutes inside the CSM continuum (i.e., the cluster-continuum model) may lead to more accurate predictions for reaction energies in solution.19,20 However, there is no general rule for choosing the number of solvent molecules and their conformational states. In the present study, we applied simple SMD model, because we focus on the relative reaction energies among different amines in which a large portion of the computational errors due to the CSM approximation would be canceled out. The Gibbs free energy (G) was calculated from the electronic energy by adding the thermal energies and entropic contributions obtained via the molecular partition functions. These calculations were performed with the default method of the Gaussian 09 package, where the translational, rotational, and vibrational partition functions were calculated based on the ideal gas, rigid rotor, and harmonic oscillator models, respectively. Calculating translational entropy in solution based on the ideal-gas model has been questioned and several correction methods have been proposed. However, it seems that there is no consensus on this issue,21 and hence, we calculated G without any corrections. We are mainly concerned with the relative G among different amines (i.e., G) in which a large portion of the computational errors due to the ideal-gas approximation would be canceled out.


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The conformational variations of an amine and its derivatives such as the carbamate anion and protonated cation may have significant effects on the thermodynamics of CO2 absorption. Although alkanolamines should be stabilized via the formation of intramolecular hydrogen bonds, a preference for such conformers depends on the competition between the intra- and intermolecular interactions. In spite of the detailed analyses of the conformational distributions,22,23 there is no consensus on this issue.24 Since the purpose of this study is to elucidate a general reaction mechanism, we have neglected the conformers that form rigid intramolecular hydrogen bonds. Within such restrictions, we have examined several conformers for every species and chose the most thermodynamically stable conformers.

RESULTS AND DISCUSSION Electronic Effects of Methyl Groups Using quantum chemical calculations, we evaluated the Gibbs free energy change, G3, for the reaction in Eq. 3 at 30 ºC. The experimental value of equilibrium constant K3 was obtained from the experimental carbamate stability constant Kc (Eq. 4) using the relation K3 = 1/(Kc[H2O]), setting [H2O] = 55.56 mol/L for infinite dilution. Following this, G3 was obtained via the relation G3 = −RTlnK3, where R is the gas constant and T is the absolute temperature. Table 1 shows that the calculated G3 values are in good agreement with the experimental values, and more importantly, the effects of the methyl groups on G3, as calculated by the relation G3 ≡ G3  G3(X) (X = MEA for alkanol amines and X = PA for alkyl amines), are well reproduced by the present calculations. It is to be noted that the experimental G3 value for AMP represents an upper limit because of the detection limit in the NMR measurements.10 A possible estimation of G3 is ca. 1.0 kcal/mol, after accounting for the corrections in the computational error in a similar amine, AP (1.28 − 0.32 = 0.96). Since our calculations could fairly capture the relevant thermodynamics, they can be safely applied to the following analyses.


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Table 1. Experimental and computational data for the stability of carbamate. Expt.

Calc.f

pKa

logKcb

G3d

G3

G1b

G1a

G2

Ga-z

MEA

9.35a

1.54a

4.56

4.32

5.97

-6.86

5.20

-6.68

AP

9.36a

0.72a

3.42 (-1.14)e

3.74 (-0.57)

5.29 (-0.68)

-5.71 (+1.15)

4.16 (-1.04)

-5.20 (+1.48)

AMP

9.54a

< -0.7a

< 1.45 (-3.11)

1.28 (-3.03)

3.77 (-2.21)

-6.55 (+0.30)

4.07 (-1.13)

−5.12 (+1.56)

4.17 (-0.14)

6.01 (+0.04)

-6.89 (-0.03)

5.05 (-0.15)

-7.22 (-0.54)

AMPb PA

10.40a

SBA

1.67a

4.74

4.33

6.38

-5.61

3.56

-5.86

1.31c

4.07 (-0.67)

3.97 (-0.36)

5.52 (-0.85)

-5.01 (+0.35)

3.46 (-0.10)

-4.83 (+1.03)

1.44 (-2.88)

4.44 (-1.94)

-5.87 (-0.26)

2.87 (-0.69)

-5.24 (+0.62)

4.79 (+0.46)

6.45 (+0.07)

-5.54 (+0.06)

3.88 (+0.33)

-6.43 (−0.57)

TBA IBA aData

10.33a

4.76 (+0.03)

from Reference 10.

bCarbamate cData

1.69a

stability constant defined by Eq. 4.

at 18 ºC, obtained from Reference 12.

dGibbs

free energy change (kcal/mol) for the reaction in Eq. 3 derived from Kc, as described in text.

eShown

in parentheses are the values of G ≡ G − G(X) (X = MEA for alkanol amines and X = PA for alkyl amines).

G at 30 ºC calculated using the CCSD(T)/6-311++G(2df,2p)//MP2/6-311++G(d,p) method, combined with the SMD

f

solvation model. Gi represents G for the reaction in Eq. i with the minus sign indicating a reverse reaction. Ga-z is the Gibbs free energy of carbamic acid relative to the zwitterion.

It should be mentioned here that similar results were also obtained in some of the previous theoretical studies,15,16,25-27 wherein it was shown that G3 < 0 for AMP. In addition, a clear correlation between G3 and the atomic charge on the amino nitrogen was demonstrated in ref 26, in line with the earlier theoretical study.13 In this report, we provide a more detailed analysis of this issue, taking into account the actual reaction pathway from the carbamate to bicarbonate. It is noteworthy that the formation of carbamate has been observed even in the case of AMP by the 13C-NMR measurement during the CO2 absorption into the 30 wt% aqueous solution.28 It has been revealed that the carbamate concentrations increase up to a loading of 0.5 mol CO2/mol AMP and then decrease with further increase in the CO2 loading.28 Moreover, several theoretical studies have investigated the formation reaction of AMP carbamate revealing that the activation energy associated with the rate-determining ACS Paragon Plus Environment

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step (Eq. 1a) is similar to that for MEA,6,16,27 and hence, the formation of AMP carbamate is kinetically probable. It has been pointed out that AMP exhibits a much higher CO2 absorption rate relative to the tertiary amine DEEA (diethylethanolamine), while both yield bicarbonate as a major product in aqueous solution, despite their similar basicity.29 We consider that the formation of carbamate, which is much faster than that of bicarbonate, might be responsible for this feature, while another mechanism, i.e., entropic effects associated with reorganization of H2O molecules adjacent to the amino nitrogen, has been proposed from the ab initio molecular dynamics study.30

O C O

HO

O

(1a)

NH2

C

O

NH2

HO

H2N

C

O

NH

OH

zwitterion O

O

(1b) HO

carbamate

H3N

OH

OH

C NH

carbamic acid

HO

O C O

O H H

H2N

(2) OH

O C OH O

H3N

bicarbonate

OH

Figure 2. Reaction pathway from carbamate to bicarbonate indicated by green arrows. Influence of the methyl groups in AMP on the equilibrium of each elementary step is shown schematically. The direction of the red arrows indicates the shift of the equilibrium relative to MEA, and the thickness of the arrow is proportional to the magnitude of G.

As mentioned above, the direct reaction in Eq. 3 is essentially prohibited, and the conversion of carbamate to bicarbonate can be established via free CO2 through the reverse reactions in Eq. 1b and 1a, followed by the reaction in Eq. 2 (Figure 2).6 Although we had proposed a direct pathway without invoking the intermediacy of free CO2 and the alkaline hydrolysis of carbamic acid constituting the key ACS Paragon Plus Environment

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step,7 the contribution of such pathway has not yet been verified. Therefore, we focused on the previously mentioned three-step pathway and calculated the G for the constituting elementary reactions, which are summarized in Table 1.

Table 2. Calculated electronic parameters for free amines and carbamates. Amine

Carbamate

QNa

nb (eV)

QN

n (eV)

QOa

MEA

-1.095

-10.649

-0.877

-10.399

-0.964

AP

-1.158 (-0.062) c

-10.594 (+0.056) c

-0.928 (-0.051)

-10.318 (+0.080)

-0.973 (-0.010)

AMP

-1.218 (-0.123)

-10.541 (+0.108)

-1.026 (-0.149)

-10.286 (+0.112)

-0.975 (-0.011)

AMPb

-1.142 (-0.047)

-10.477 (+0.172)

-0.819 (+0.058)

-10.227 (+0.171)

-0.942 (+0.022)

PA

-1.158

-10.546

-0.900

-10.280

-0.943

SBA

-1.166 (-0.007)

-10.588 (-0.042)

-0.927 (-0.027)

-10.222 (+0.058)

-0.954 (-0.011)

TBA

-1.271 (-0.113)

-10.587 (-0.040)

-1.068 (-0.169)

-10.271 (+0.010)

-0.977 (-0.033)

IBA

-1.172 (-0.014)

-10.451 (+0.096)

-0.892 (+0.008)

-10.212 (+0.068)

-0.955 (-0.012)

aAtomic

charges on the amino nitrogen atoms (QN) and OCO oxygen atoms (QO) as calculated using the ChelpG scheme

with MP2 density. bHartree-Fock

cShown

orbital energy of the nitrogen lone-pair orbitals.

in parentheses are the values relative to MEA for alkanol amines and PA for alkyl amines.

Since the variation in G2 arises from the difference in the basicity of amines, our calculations suggest that all amines have higher basicity than MEA, which is in agreement with the trends in the experimental pKa values (Table 1). It should be noted that in the present study, the translational entropy was calculated using the ideal gas formula that significantly overestimates the entropy of every species in solution. Thus, for the reactions in which the number of molecules is changed, e.g., the reactions in Eq. 1a and 2, only G are meaningful and are sufficient for discussion. The atomic charges on the nitrogen atoms calculated using the ChelpG scheme31 based on the MP2 density, together with the ACS Paragon Plus Environment

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Hartree-Fock orbital energy of nitrogen lone pair orbitals, are listed in Table 2 for both free amines and carbamates. We note that the higher basicity (larger pKa) of the amino nitrogen is associated with the larger negative charge and higher lone pair orbital energy on the nitrogen atom, as can be intuitively expected from the electron-donating nature of methyl groups. Our calculations predict that the significantly reduced thermodynamic stability (decreased G3) of the AMP carbamate, relative to that of MEA, can be largely ascribed to the preferred reverse reaction in Eq. 1b (decreased G1b). Since the equilibrium of the reaction in Eq. 1b is governed by the basicity of the amino nitrogen in carbamate relative to that in free amine, the decreased G1b implies that the increase in the basicity of nitrogen on going from MEA to AMP is more significant in carbamate than in free amine. This behavior is likely to be caused by the electronic effects of the methyl substituents, because it is consistent with the trends in the atomic charges and the orbital energies (Table 2). The increase in the negative charge on the nitrogen atom of AMP, relative to that of MEA, is more significant in carbamate (0.149) than in free amine (0.123), and the upward shift of the nitrogen lone pair orbital energy in carbamate (+0.112 eV) is slightly larger than that in free amine (+0.108 eV). However, the steric effects of the methyl groups would be less significant because the calculated C-N bond length in AMP carbamate (1.410 Å) is almost identical to that of MEA carbamate (1.409 Å). The reverse reaction in Eq 1a represents the release of CO2 from the zwitterion. The positive G1a of AMP (Table 1) implies higher stability of the AMP zwitterion (i.e., larger binding energy of CO2 to amine) compared with that of MEA. In this regard, it is interesting to note that a much larger stabilization of the zwitterion is obtained via mono substitution, as evident from the comparison of AP and AMP, as well as SBA and TBA. The protonation of carbamate at the oxygen atom of OCO moiety results in the formation of carbamic acid, which competes with the formation of zwitterion (the reverse reaction in Eq. 1b). The carbamic acid acts as a trap against the decomposition of carbamate, because it has no direct channel to release CO2. As a result, an increasing concentration of the carbamic acid, relative to the zwitterion, ACS Paragon Plus Environment

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would reduce the rate of carbamate decomposition. The calculated Gibbs free energies of the carbamic acid relative to the zwitterion, Ga-z, are listed in Table 1. Although carbamic acid is more stable than the zwitterion for all the amines investigated, the relative stability of the former is reduced in AMP, thus increasing the rate of transformation from carbamate to bicarbonate. In terms of the electronic properties, the increase in the basicity of the carbamate on going from MEA to AMP is more significant for the amino nitrogen than for the OCO oxygen, as indicated by the ChelpG atomic charges shown in Table 2. It is evident that the increase in the negative charge on the amino nitrogen (0.149) is much larger than that on the OCO oxygen (0.011). The influence of methyl groups in AMP on the equilibrium of each elementary step is schematically shown in Figure 2. Both the thermodynamic and kinetic preferences for the bicarbonate over the carbamate in AMP would be specific to the substitution of methyl groups near the amino nitrogen. This is because the same substitutions on the -carbon (AMPb in Figure 1) impart less significant effects on G3 and opposite effects on Ga-z. The calculated G3 for AMPb implies less significant reduction of the thermodynamic stability of carbamate compared with AMP. The G1b of AMPb is almost unchanged compared with that of MEA, indicating that the basicity of both free amine and carbamate are almost unaffected by the methyl groups on the -carbon. On the other hand, AMPb shows increased thermodynamic stability of the carbamic acid relative to the zwitterion, in contrast to that observed for AMP. All these facts are consistent with the trends in the ChelpG atomic charges listed in Table 2. A similar site-dependence of the electronic effects of the methyl groups is also found in the case of alkyl amines (PA and its derivatives, Figure 1). Both experiments and our calculations show that the methyl substitutions at the -carbon of PA (giving SBA and TBA) reduce the thermodynamic stability of carbamates (Table 1), with similar G3 values for the corresponding alkanol amines (AP and AMP). In contrast, the methyl substitution on the -carbon (giving IBA) does not reduce the stability of carbamate, as in the case of the alkanol amine (AMPb). In addition, the influence of substitution sites on the Ga-z is similar to that of alkanol amines, and all these thermodynamic features in alkyl amines are ACS Paragon Plus Environment

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consistent with the trends in the atomic charges. It should be mentioned here that the entropic effects would not be relevant for all the substituent effects described above, because the Gi (i = 3, 1b, 1a, 2, a-z) are quite close to the corresponding electronic energies, Ei, as shown in Table S1.

Steric Effects of Methyl Groups The influence of steric hindrance from the substituent components is investigated in the context of conformational dynamics. We focus on the rotation of the OCO moiety that corresponds to the lowest frequency normal vibrational mode of the zwitterion according to our calculations, for both MEA and AMP. Figure 3 shows the potential energy profiles against the O1-C2-N3-C4 torsional angle for both the zwitterion and carbamate, as obtained by the optimization of all the remaining geometrical parameters. In both MEA and AMP, the rotation of OCO moiety is associated with a rather low energy barrier (2-3 kcal/mol) in the zwitterion, while it is almost prohibited in the carbamate. Therefore, at operating temperatures, this feature would allow the zwitterion to adopt a conformation in which CO2 is forced out from the AMP zwitterion due to the steric repulsion imposed by the methyl groups.

15

15

(b) AMP

10

carbamate 5

zwitterion 0

Energy (kcal/mol)

(a) MEA Energy (kcal/mol)

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


O1

10

H3C

carbamate 5

O

C4 N3

HO

zwitterion

C2

H

C5 H6 H

0 0

50

100

150

Torsion Angle (degree)

0

50

100

150

Torsion Angle (degree)

Figure 3. Potential energies calculated using SMD-MP2/6-311++G(d,p) with respect to the O1-C2N3-C4 torsion angle for the zwitterion and carbamate of (a) MEA and (b) AMP.


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An example of such zwitterion conformer is illustrated in Figure 4(b); the most stable conformer is also provided for comparison (Figure 4(a)). The former is characterized by the torsional angles of O1C2-N3-C4 = C5-C4-N3-C2 = H6-C5-C4-N3 = 0º. In Figure 4(b) the distance between O1 and H6 is 1.435 Å, which is much shorter than the sum of the van der Waals radii of oxygen (1.4 Å) and hydrogen (1.2 Å). Since it requires only 7.8 kcal/mol to convert the latter two torsional angles to 0º, the AMP zwitterion may instantaneously adopt such a conformation, with CO2 being forced out from AMP due to the steric repulsion. Recent ab initio molecular dynamics simulations have revealed that the release of CO2 from MEA carbamate proceeds via the zwitterion intermediate.32 It might be interesting to examine whether such an event takes place in AMP. The regeneration process of CO2 from carbamate can be governed by the relative rates between CO2 release (reverse reaction in Eq. 1a) and deprotonation (reaction in Eq. 1b) from the zwitterion. Several theoretical studies based on the stable conformers have predicted that the activation energy for the former reaction is larger than that for the latter with no significant difference between MEA and AMP.6,7,16,27 The present study suggests a new scenario unique to AMP in which the methyl groups promote the release of CO2 from the zwitterion.

(a)

(b) 1.435 Å

Figure 4. SMD-MP2/6-311++G(d,p) calculated structures of the AMP zwitterion with (a) minimum energy and (b) hindered conformation.

CONCLUSIONS ACS Paragon Plus Environment

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Using ab initio calculations combined with the SMD solvation model, we have investigated the origin of the reduced stability of AMP carbamate from both thermodynamic and kinetic points of view. The calculated Gibbs free energy changes (G3) for the formal hydrolysis of carbamate to bicarbonate are in sufficient agreement with those obtained from the NMR speciation. We have concluded that the reduced thermodynamic stability of AMP carbamate is largely ascribable to the shift of the equilibrium toward the zwitterion in Eq 1b. The shift is caused by the electronic effects of the methyl groups that increase the basicity of amino nitrogen more significantly in carbamate than in free amine. In addition, the methyl groups in AMP enhance the stability of the zwitterion relative to carbamic acid. Since there is no direct route for the release of CO2 from carbamic acid, this feature increases the rate of CO2 release from the carbamate. We further revealed that these thermodynamic and kinetic features of AMP is specific to the attachment of methyl groups on the -carbon, because the same substitutions at the -carbon imparts less significant effects on the equilibrium of Eq. 1b and opposite effects on the equilibrium between the zwitterion and carbamic acid. Such site-dependent substitution effects were also confirmed to hold true for alkyl amines. The steric effects of methyl groups on the stability of AMP carbamate have been investigated taking into account the conformational dynamics of zwitterion intermediate. We found that the rather small energy barrier for the rotation of the OCO moiety in the zwitterion allowed it to adopt a conformation that had significant steric repulsion between one of the methyl groups and the OCO moiety, which may lead to the release of CO2 from the zwitterion intermediate, thus promoting the decomposition of carbamate.

Acknowledgment. This work was financially supported by the COURSE50 project founded by the New Energy and Industrial Technology Development Organization, Japan.


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Supporting Information Available: Table S1 and complete Reference 17. This material is available free of charge via the Internet at http://pubs.acs.org.

References and Notes (1) Kerr, R. A. Global Warming Is Changing the World. Science 2007, 316, 188-190. (2) Rochelle, G. T. Amine Scrubbing for CO2 Capture Science 2009, 325, 1652-1654. (3) Caplow, M. Kinetics of Carbamate Formation and Breakdown. J. Am. Chem. Soc. 1968, 90, 67956803. (4) Danckwerts, P. V. The Reaction of CO2 with Ethanolamines. Chem. Eng. Sci. 1979, 34, 443-446. (5) Xie, H-B.; Zhou, Y.; Zhang, Y.; Johnson, J. K. Reaction Mechanism of Monoethanolamine with CO2 in Aqueous Solution from Molecular Modeling. J. Phys. Chem. A 2010, 114, 11844-11852. (6) Yamada, H.; Matsuzaki, Y.; Higashii, T.; Kazama, S. Density Functional Theory Study on Carbon Dioxide Absorption into Aqueous Solutions of 2-Amino-2-methyl-1-propanol Using a Continuum Solvation Model. J. Phys. Chem. A 2011, 115, 3079-3086. (7) Matsuzaki, Y.; Yamada, H.; Chowdhury, F. A.; Higashii, T.; Onoda, M. Ab Initio Study of CO2 Capture Mechanisms in Aqueous Monoethanolamine: Reaction Pathways for the Direct Interconversion of Carbamate and Bicarbonate. J. Phys. Chem. A 2013, 117, 9274-9281. (8) Sartori, G.; Savage, D. W. Sterically Hindered Amines for CO2 Removal from Gases. Ind. Eng.

Chem. Fundam. 1983, 22, 239-249. (9) Yamada, H.; Shimizu, S.; Okabe, H.; Matsuzaki, Y.; Chowdhury, F. A.; Fujioka Y. Prediction of the Basicity of Aqueous Amine Solutions and the Species Distribution in the Amine-H2O-CO2 System Using the COSMO-RS Method. Ind. Eng. Chem. Res. 2010, 49, 2449-2455.


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Table of Contents Graphic

Methyl substituents shift the equilibrium O C O

HO

O

(1a)

NH2

C

O

NH2

HO

zwitterion O

O

(1b) H2N

C

O

NH

OH

HO

H3N

carbamate

OH

OH

C NH

carbamic acid

HO

O C O

O H H

H2N

(2) OH

O C OH O

bicarbonate

H3N

OH


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Ab Initio Study of CO2 Capture Mechanisms in Aqueous 2-Amino-2

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