CHEMPHYSCHEM ARTICLES DOI: 10.1002/cphc.201300820
Carbamate Stabilities of Sterically Hindered Amines from Quantum Chemical Methods: Relevance for CO2 Capture Satesh Gangarapu,[a] Antonius T. M. Marcelis,[a] and Han Zuilhof*[a, b] The influence of electronic and steric effects on the stabilities of carbamates formed from the reaction of CO2 with a wide range of alkanolamines was investigated by quantum chemical methods. For the calculations, B3LYP, M11-L, MP2, and spincomponent-scaled MP2 (SCS-MP2) methods were used, coupled with SMD and SM8 solvation models. A reduction in carbamate stability leads to an increased CO2 absorption capacity of the amine and a reduction of the energy required for sol-
vent regeneration. Important factors for the reduction of the carbamate stability were an increase in steric hindrance around the nitrogen atom, charge on the N atom and intramolecular hydrogen bond strength. The present study indicates that secondary ethanolamines with sterically hindering groups near the N atom show significant potential as candidates for industrial CO2-capture solvents.
1. Introduction The increased emission of anthropogenic greenhouse gases (GHGs) has the potential to significantly increase the average global temperature.[1] Carbon dioxide is considered one of the major contributors to global warming, due to its abundance compared to other greenhouse gases.[2] Therefore, reduction of the atmospheric emission of CO2 is considered a primary target in environmental policies. Absorption of CO2 with alkanolamines in mixtures with absorption promoters has traditionally been used for its removal from exhaust gases. Among the available alkanolamines, monoethanolamine (MEA) is the most commonly used industrial solvent for CO2 capture.[3] This is because MEA has a high reactivity towards CO2, low solvent cost, low molecular weight and good water solubility.[4] However, MEA has certain disadvantages, such as the formation of rather stable carbamates with CO2 and a concomitantly high energy demand for regeneration, a relatively high volatility, and it shows significant long-term thermal and oxidative degradation. These factors make MEA a suboptimal solvent for CO2 capture from power plants.[5] Clearly, there is considerable scope for improving CO2-capturing solvents, which have higher CO2 absorption rates, lower regeneration costs and lower degradation rates than MEA. Recently, sterically hindered primary amines and secondary amines have received much attention as CO2-capturing sol[a] S. Gangarapu, Dr. A. T. M. Marcelis, Prof. H. Zuilhof Laboratory of Organic Chemistry Wageningen University Dreijenplein 8 6703 HB Wageningen (The Netherlands) Tel: (+ 31) 317-482361 E-mail: [email protected] [b] Prof. H. Zuilhof Department of Chemical and Materials Engineering King Abdulaziz University, Jeddah (Saudi Arabia) Supporting Information for this article is available on the WWW under http://dx.doi.org/10.1002/cphc.201300820.
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vents due to their fast reaction rates with CO2 and lower regeneration costs than conventional primary amines like MEA.[6] Such sterically hindered amines form less stable carbamates upon reaction of CO2 and thereby increase the CO2 absorption capacity from 0.5 mole of CO2 per mole of amine to up to 1 mole of CO2 per mole of amine (Scheme 1). Widely used ster-
Scheme 1. Equations describing the possible reactions of alkyl amines and carbamates with CO2 and water.
ically hindered amines for CO2 capture are 2-amino-2-methyl-1propanol (AMP) and diethanolamine (DEA). Despite the advantages provided by their bulkier substituents, these solvents are not ideal due to a poor reactivity towards CO2 and improvements are still needed. In the search for new solvents for CO2 capture and optimal use thereof, we therefore need an improved and more detailed understanding of the reaction of amines with CO2, including an accurate estimate of the factors that determine their corresponding carbamate stability. In the present study, we thus investigate the electronic and steric effects on the carbamate stabilities of four series of alkanolamines, and we do this by calculating the Gibbs free energy for Reaction (C) with quantum chemical methods. The first series are primary alkanolamines in which one or two CH3 groups are present at the a or b position from the amine group (Figure 1: entries 1–5). The second series consists of secondary alkanolamines that have an increasingly bulky alkyl ChemPhysChem 2013, 14, 3936 – 3943
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www.chemphyschem.org and [Eq. (3)]: DDGs ¼ DGs ðRNH2 Þ þ DGs ðHCO3 ÞDGs ðRNHCO2 ÞDGs ðH2 OÞ
Figure 1. Alkanolamines for CO2 capture as investigated in this study. Different series of alkanolamines are separated by dotted lines.
group on the N atom (Figure 1: entries 6–12). To the best of our knowledge, the effects of bulky groups like isopropyl, isobutyl or secondary butyl substituents on the N atom of MEA (1), and N-methyl or N-ethyl substitution on AMP (5) for application in CO2 capture were not studied yet. The third series are secondary alkanolamines with various N-bound substituents like hydroxyalkyl or aminoalkyl with variable chain lengths are present (Figure 1: entries 13–17). Finally, the fourth series consists of the cyclic amines 2-piperidinemethanol and 2-piperidineethanol (Figure 1: entries 18, 19), which have already shown some promise in CO2 capture.[4c] To this aim, B3LYP/6-311 + G(2d,2p), MP2/6-311 + G(2d,2p), spin-component-scaled MP2/6-311 + G(2d,2p) (SCS-MP2)[7] and M11 L/6-31 + G(d,p)[8] calculations were performed to obtain the Gibbs free energy for the carbamate hydrolysis reaction. Solvents effects were mimicked with the universal solvation models SM8 and SMD.[9] Through this approach, we aim to delineate the electronic and steric effects on the carbamate stabilities of sterically hindered amines. Based on these finding we propose some interesting alkanolamines for industrial use as CO2 capture agents.
Computational Methods and Procedures The carbamate stabilities towards hydrolysis were calculated by using the thermodynamic cycle reported by da Silva et al.[10] The reaction energy in solution (DGcs) can be written as Equation (1): DGcs ¼ DGcg þ DDGs where [Eq (2)]: DGcg ¼ Gg ðRNH2 Þ þ Gg ðHCO3 ÞGg ðRNHCO2 ÞGg ðH2 OÞ
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ð1Þ
Gas-phase conformational analysis was done on the species that are described in the thermodynamic cycle to find the lowestenergy conformers. A conformer search was performed with the MMFF94[11] force field in ChemBio3D.[12] The lowest-energy conformers in the solution phase were found to be the same as in the gas phase. The pre-optimized gas-phase geometries obtained with MMFF94 were further optimized in solution phase with B3LYP/6311 + G(2d,2p), MP2/6-311 + G(2d,2p) and M11-L/6-31 + G(d,p) calculations. Frequency analyses were carried out on the optimized geometries with the same basis set to find true minima on the potential energy surface. Thermal corrections and entropic contributions were added to the total electronic energy for the calculation of the free energy. The universal solvation models (SMD and SM8) that we used for the calculation of the free energy of solvation are trained against the experimental free energies of solvation, and include an entropic contribution to the free energy of solvation. All calculations were carried out with the Gaussian 09[13] software package, Q-chem 4.0[14] and Gaussview 5.0 visualization software. B3LYP, MP2, SCS-MP2 calculations were carried out using Gaussian 09 with the use of the SMD solvent model with the 6-311 + G(2d,2p) basis set.[15] In case of M11L, calculations were carried out using Q-Chem 4.0 [which allows the use of the SM8 solvent model with basis sets up to 6-31 + G(d,p)]. Recently, it was shown that SCS-MP2 methods give significant improvements in cases where MP2 underperforms, which prompted us to use SCS-MP2 energies wherever possible.[7, 16] All data in the tables below represent SMD/ B3LYP/6-311 + G(2d,2p)//SMD/B3LYP/6-311 + G(2d,2p), SMD/SCSMP2/6-311 + G(2d,2p)//SMD/MP2/6-311 + G(2d,2p) calculations. Data of analogous SM8/M11-L/6-31 + G(d,p)//SM8/M11-L/6-31 + G(d,p) calculations are given in the Supporting Information. Atomic charges of all species were calculated by natural population analysis (NPA)[16b] and hybridizations by NBO analysis (NBO)[16b] on the solution phase-optimized geometries.
2. Results and Discussion The geometries of the alkanolamines and their carbamates were optimized with SCS-MP2/6-311 + G(2d,2p) and M11-L/631 + G(d,p) methods. From previous studies, it was found that both methods yield highly accurate pKa and carbamate stability data and thus display the same trends.[17] The lowest-energy conformers of carbamate anions of selected alkanolamines after MP2-optimization, along with their intramolecular hydrogen bond lengths and NC bond lengths are shown in Figure 3. All other MP2-optimized geometries, together with the M11-optimized geometries of all compounds under study, are shown in the Supporting Information (Figures S1 and S2, respectively). The lowest-energy conformers of the carbamate anions show H-bonds between an O atom of the COO group and the H atom of the OH group. The lowest-energy conformers of the alkanolamines show intramolecular H-bonding between the amine N atom and the H atom of the hydroxyl group. The lowest-energy conformers of either alkanolamine or its carbamate anion in the gas phase and solution phase are almost identical in geometry, and the same intramolecular Hbonds are observed in the carbamate anions in the gas phase and in solution.[18] The Gibbs free energy as calculated from ChemPhysChem 2013, 14, 3936 – 3943
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Figure 2. Thermodynamic cycle used for the calculation of the Gibbs free energy of Reaction (C).
the thermodynamic cycle (Figure 2) was used as a parameter to interpret the carbamate stabilities of the sterically hindered amines.
2.1. Substitution of C-linked Methyl Groups on MEA For the sterically hindered primary amines, the effect of substitution of one or two methyl groups either on the a- or bcarbon atom from the amine functional group was investigated. The lowest-energy conformers of the carbamate anions of relatively important alkanolamines are shown in Figure 3, to-
Figure 3. Lowest-energy conformers of the carbamate anions formed upon reaction of CO2 with some relatively important substituted alkanolamines, showing also the O···HO and NC bond lengths [SCS-MP2/6-311 + G(2d,2p) data]. Black: N; dark grey: O; grey: C; white: H.
gether with their H-bond and NC bond lengths. The carbamate stability of aCH3 substituted MEA (3) was found to be lower than the carbamate stability of MEA (1), due to the increase in steric hindrance around the amine group (Table 1). Further substitution
with a second methyl group on the a-position to the amine group yields the well-known alkanolamine AMP (5), which displays a further decreased carbamate stability of 3.59 kcal mol1. The reduction in the intramolecular hydrogen bond length between the hydrogen atom from the hydroxyl group and the oxygen atom from the COO group of a carbamate with one or more CH3 groups on the a-carbon atom is reflected by a decrease in NCC(OH) bond angle, which decreases stepwise from 113.48 via 111.98 to 111.38 in going from MEA-carbamate (1) via the carbamate of 3 to that of 5. However, the N CC(OH) bond angle of 2 is 113.78, which is similar to that of the MEA carbamate (1), indicating that there is almost no effect of the bond angle due to substitution of a CH3 group on the b-carbon atom from the amine group. The decrease in the NCC(OH) bond angle upon substitution of CH3 groups on the a-carbon atom is due to the so-called Thorpe–Ingold effect.[19] A result of a smaller NCC(OH) angle is that the intramolecular (OC)O···H(O) bond becomes shorter. The increase in steric hindrance thus corresponds with a decrease in the intramolecular H-bond length, with (OC)O···H(O) distances for the carbamate of 1 of 1.630 to 1.612 for 3 and 1.589 for 5. The effect of intramolecular hydrogen bonding in the a-CH3 substituted alkanolamines is further studied by looking at the structural changes between the carbamate of 5 with and without intramolecular hydrogen bonding. The NCC(OH) bond angle of 5 with hydrogen bonding was found to be 111.38, which is higher than the NCC(OH) bond angle of 106.78 in the non H-bonded carbamate. The increase in the NCC(OH) bond angle of 5 with hydrogen bonding is accompanied by a decrease in CH3CCH3 bond angle from 109.38 to 107.98. The decrease in the carbamate stability of a-CH3 substituted alkanolamines can be further explained by looking at the NC bond lengths in the carbamate species of substituted MEA. The NC bond length in the carbamate of unsubstituted MEA is 1.443 , which is characteristic for an NC single bond.[20] Upon substitution of a -CH3 groups on MEA, the NC bond length increases (data shown in supporting information Table S2), in line with a weakening of the NC bond by substitution of hydrogen by a -CH3 group. This weakening is crucial for a more cost-effective regenerative desorption of CO2. Both SMD/SCS-MP2 and SM8/M11-L calculations show that DG for the carbamate hydrolysis reaction becomes less negative with an increase in steric hindrance around the N atom. The carbamate stability towards hydrolysis of substituted alkanolamines follows the order: MEA > 1-amino-2-propanol > 2amino-1-propanol > 3-amino-1-propanol > AMP, which follows
Table 1. Alkyl-substituted MEAs: DGcs, NCC bond angle, OH bond lengths and NPA charges on the N atoms. Compound Entry
DGcs [kcal mol1]
ffNCC [8]
r(OH) []
NPA charge on N atom
monoethanolamine (MEA) (1) 1-amino-2-propanol (2) 2-amino-1-propanol (3) 3-amino-1-propanol (4) 2-amino-2-methyl-1-propanol (AMP) (5)
5.83 5.34 5.04 4.96 3.59
113.4 113.7 111.9 – 111.3
1.630 1.618 1.612 1.640 1.589
0.882 0.879 0.879 0.877 0.865
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the same trend as experimentally found for compounds for which these data are available.[4c] The calculated DGcs of MEA carbamate was 5.83 kcal mol1, 6.79 kcal mol1, 6.28 kcal mol1 and 6.55 kcal mol1 with SCS-MP2, B3LYP, MP2 and M11-L respectively, that is, in good agreement with the experimentally reported value of 4.92 kcal mol1 for the reverse reaction.[21] Although accurate experimental data are scarce, this suggests that the methods used (SMD/SCS-MP2, SMD/MP2, SMD/B3LYP and SM8/M11-L) can accurately calculate the Gibbs free energy for the carbamate hydrolysis reaction. To check for dependence of the overall result on the solvent model and basis set, we performed the calculation on three compounds from different series that is, 1, 8, and 18, with SMD/M11-L/6-311 + G(2d,2p) (data shown in Table S1) and compared it with one used in the calculations that is, SM8/M11-L/6-31 + G(d,p). It was found that the results from both models are almost identical. The DGcs values from B3LYP, M11 L and MP2 deviate from the SCS-MP2 values by ~ 1.2 kcal mol1, ~ 1.0 kcal mol1 and ~ 0.5 kcal mol1, respectively. However, the results from SCSMP2 are more accurate than those from M11-L, and therefore the results from SCS-MP2 are discussed here and data from B3LYP, MP2 and M11-L are shown in Tables S1 and S2. A good correlation (R2 = 0.96) was observed between DGcs for the carbamate hydrolysis reaction and the charge on the N atom in alkanolamines (Figure 4 a). It clearly indicates that the higher the charge on the N atom, the stronger the bond to CO2. A very good correlation (R2 = 0.99) was observed between DGcs and the O···HO intramolecular hydrogen bond length in carbamates (Figure 4 b). 2.2. Substitution of N-Linked Alkyl Groups on MEA In this section, the direct substitution of an H atom on the amine N atom in MEA by a variety of increasingly bulky alkyl groups is investigated. The DGcs, intramolecular H-bond lengths, NCC bond angles and NPA charge on the N atom of the carbamate anions of N-alkyl-substituted MEAs are shown in Table 2. As expected, direct substitution of a CH3 group on the N atom of the amine functional group (6) decreases the carbamate stability of the alkanolamine, as seen from the decrease in the Gibbs free energy of reaction from 5.83 kcal mol1 to
Figure 4. Correlation plots between DGcs for the carbamate hydrolysis reaction and a) the NPA charge on the N atom, or b) DGcs the O···HO intramolecular H-bond length in the carbamates of analogous compounds (1–3 and 5) (4 is excluded from the graph, because it is an aminopropanol).
4.80 kcal mol1 (Table 2). Substitution by an ethyl group (7) on the N atom yielded a DGcs of 4.53 kcal mol1 (Table 2), which is lower than for the alkanolamines 1 and 6. This is obviously due to the inductive effect of a CH3 group on the N atom. The total NPA charge of the CH3 group in the carbamate of 6 is 0.357, which is more positive than the NPA charge on the hydrogen atom (0.206) in carbamate 1. Further increase in steric hindrance around the N atom by substitution with an isopropyl group yielded the desired clear decrease in carbamate stability. The DGcs of the isopropyl group substituted alkanolamine 8 was found to be ~ 4.3 kcal mol1 smaller (i.e., less negative) than of unsubstituted MEA (1): only 1.56 kcal mol1 (Table 2). Surprisingly, substitution by more bulky groups on the N atom, like isobutyl (9) or secondary butyl (10), resulted in a more negative Gibbs free energy than obtained for the isopropyl group.
Table 2. N-substituted MEAs and 2-aminopropanols: DGcs, NCC bond angle, OH bond lengths and NPA charges on the N atoms. Compound Entry
DGcs [kcal mol1]
ffNCC [8]
r(OH) []
NPA charge on N atom
2-(methylamino)ethanol (6) 2-(ethylamino)ethanol (7) 2-(1-methyl)ethylamino) ethanol (8) 2-(2-methyl)propylamino) ethanol (9) 2-(1-methylpropylamino) ethanol (10) 2-(N-methylamino)-2-methyl-1-propanol (11) 2-(N-ethylamino)-2-methyl1-propanol (12)
4.80 4.53 1.56
113.6 113.9 114.1
1.622 1.630 1.646
0.738 0.736 0.729
2.85
114.0
1.642
0.731
3.53
114.0
1.638
0.735
1.41 0.84
112.2 111.6
1.575 1.568
0.744 0.740
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CHEMPHYSCHEM ARTICLES We hypothesized that this might be due to the difficulty of solvent accessibility during the solvent energy calculations with the bulky groups. To test this idea, the solvation energies for alkanolamines 6–10 were calculated with inclusion of one explicit water molecule in the solvation energy calculations. The results of including one explicit water molecule are shown in Figure 5. The distance between the N atom of the carba-
Figure 5. Optimized structures obtained by inclusion of an explicit water molecule to the reaction center (N atom) in alkanolamines 6–10. Black: N; dark grey: O; grey: C; white: H.
mate and O atom of this solvating water molecule is 2.334 for 6, 2.626 for 7 and 3.192 for 8. For alkanolamine 10, which contains the bulky secondary butyl group, the distance is even > 4 . This clearly indicates that upon increasing the steric hindrance around the nitrogen atom in alkanolamines, it becomes more difficult for the solvent to interact with the reaction center. This observation is further supported by calculating the solvation energies of carbamates of 6–10. The solvation energies for secondary butyl (10) and isobutyl (9) substituted MEA are ca. 2–3 kcal mol1 smaller than the solvation energy for the sterically less hindered compounds 6–8 (full data shown in Table S4). Therefore, the steric hindrance around the N atom seems to have an optimum when aiming for an improved CO2 absorption capacity, that is, with the isopropyl group. By comparing DGcs of the carbamates of ethanolamines with different substituents on the N atom, it is seen that substitution of a H atom by either a methyl, ethyl or isopropyl group improves the CO2 absorption capacity by a decrease in the stability of the carbamate anions. The changes in the intramolecular H-bond length in N-alkyl MEA nicely correspond with the changes in NCC angles. Substitution of an ethyl group (7) on the N atom increased ffNCC from 113.48 for MEA to 113.98 for 7, and increased the intramolecular H-bond length from 1.620 to 1.630 . Further substitution of an isopropyl group on MEA (8) increased ffNCC to 114.18 and the intramolecular H-bond length to 1.646 .
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www.chemphyschem.org A good correlation (R2 = 0.94) was observed between DGcs and the charge on the N atom in the ethanolamines (Figure 6 a). This also clearly indicates that the higher the charge
Figure 6. Correlation plots between a) the Gibbs free energy for the carbamate hydrolysis reaction (DGcs) and NPA charge on N atom, and b) DGcs and O···HO intramolecular hydrogen bond length in carbamates of ethanolamines 6–10.
on the N atom, the stronger its binding to CO2. A reasonably good correlation (R2 = 0.89) was observed between DGcs and the O···HO intramolecular H-bond length in carbamates (Figure 6 b). It was found that a longer H-bond length, that is, weakened H-bond in the reactant, decreases the resulting carbamate stability, which is in contrast with the observation mentioned in Section 2.2. This difference is due to the direct substitution of alkyl groups on the N atom in these molecules. Due to the increasing steric hindrance, the bond angles around the N atoms increase to close to 1208 in compounds 9 and 10, corresponding to a hybridization change from sp3 to almost sp2.[22] The change in the hybridization around the N atom was further confirmed by NBO analysis of the carbamates of 9 and 10. The correlation plot between DGcs and the charge on the nitrogen atom in all alkanolamines included in this study is shown in Figure S3, which shows that DGcs and the NPA charge depend strongly on the nature of the amine group in alkanolamines. Methyl substitution on the N atom of AMP (11) decreased the Gibbs free energy of reaction from 3.59 kcal mol1 to 1.41 kcal mol1 (Table 2). Furthermore, substitution of an ethyl group on AMP (12) decreased the Gibbs free energy of reaction to 0.84 kcal mol1, also due to poor solvation (Table 2). This clearly indicates that upon increase in the steric hindrance around the N in the alkanolamine, the carbamate stability decreases, thereby increasing the CO2 absorption capacity of the ChemPhysChem 2013, 14, 3936 – 3943
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alkanolamine. This means that substitution of one of the hydrogen atoms on the N of AMP by either a methyl or ethyl group provides a highly promising route to reduce the energy needed for regeneration of alkanolamines from carbamates.
2.3. Substitution of Hydroxyalkyl and Aminoalkyl Functional Groups on the N Atom in MEA In this section, the carbamate stabilities of a series of alkanolamines that contain alkylhydroxy or alkylamine functional groups (13–17) were calculated and compared with MEA (1). The DGcs, intramolecular H-bond lengths, ffNCC and NPA charge on the N atom of carbamate anions of all alkanolamines (13–17) are shown in Table 3. Upon substitution of an N-bound H atom by an N-hydroxyethyl (CH2CH2OH) group (13) or an aminoethyl (CH2 CH2NH2) group (14), DGcs decreased from 5.83 kcal mol1 to 4.76 and 4.51 kcal mol1, respectively. This decrease in carbamate stability is attributed to the electron-withdrawing nature of these groups. The formation of less stable carbamates from 13 and 14 is further supported by the NPA charge on the nitrogen atom. The NPA charges on the N atoms in both 13 (0.749) and 14 (0.734), are smaller than the value of 0.882 found for MEA (1), indicating a reduced reactivity towards CO2. Compounds 14 and 17 contain two reaction centers that is, both primary and secondary amine groups. We considered the formation of a carbamate on the secondary amine group, due its higher reactivity towards CO2 than a primary amine group. In compound 15, in which two 2-hydroxypropyl moieties are present on the N atom, the carbamate stability decreases due to the increase in the steric hindrance. Compound 16 forms the least stable carbamate among all amines in Table 3, despite having a primary amine group. This is caused by the formation of a second H-bond (OH···N) in this compound. This is further indicated by the longest NC bond length (data shown in the Supporting Information) in the carbamate anion of compound
16. A very poor correlation was observed between DGcs either with the NPA charge on the N atom or with the intramolecular H-bond length. This may be due to the presence of extra intramolecular H-bonding in the alkanolamines 13–17. Data from both calculations (SCS-MP2 and M11-L) show that the carbamate stabilities decrease in the following order: MEA > DEA > AMP, which is the same as the experimental values reported by Satori and Savage.[4c] It clearly confirms that both methods (SMD/SCS-MP2 and SM8/M11-L) are good enough to calculate the carbamate stabilities of alkanolamines accurately.
2.4. Linear Alkanolamines versus Cyclic Alkanolamines In this section, the carbamate stabilities of linear and cyclic alkanolamines are investigated. Cyclic alkanolamines were found to exhibit higher CO2 absorption capacities than conventional alkanolamines.[4c, 23] Table 4 shows the DGcs, H-bond lengths, NCC bond angles and the NPA charge on the N atom of the carbamate anions of cyclic alkanolamines 18 and 19 and of the corresponding linear forms 1 and 4. It was found that DGcs of the carbamates of the linear alkanolamines are higher than of the cyclic alkanolamines with the same number of carbon atoms between the hydroxyl and amine groups. Upon increasing the carbon chain length between the amine and hydroxyl functional groups, the carbamate stability decreased in both linear and cyclic alkanolamines. This is caused by a reduction of the electron withdrawing effect of the hydroxyl group on the nitrogen atom of the amine.[24] In conclusion, the carbamate stability of substituted alkanolamines does not only depend on the different electron-donating or -withdrawing properties of the functional groups, but also depends on the position (a or b or direct) of a substituent with respect to the amine or hydroxyl functional group or linearity of the molecule or the carbon chain length between the amine and hydroxyl functional groups in the alkanolamines. From the calculations we can conclude that upon increasing
Table 3. Hydroxyalkyl- and aminoalkyl substituted MEAs: DGcs, NCC bond angle and OH bond lengths, and NPA charges on the N atoms. Compound Entry
DGcs [kcal mol1]
ffNCC [8]
r(OH) []
NPA charge on N atom
di-(2-hydroxyethyl) amine (13) 2-(2-aminoethyl)amino ethanol (14) 1-(2-hydroxypropyl) aminopropanol (15) 2-amino-2-methyl-1,3-propanediol (16) 2-(3-aminopropyl) aminoethanol (17)
4.76
114.1
1.689
0.749
4.51
113.9
1.633
0.734
4.23 1.90 4.49
114.6 114.2 114.1
1.668 1.674 1.643
0.744 0.882 0.729
Table 4. Linear alkanolamines 1 and 4 and cyclic alkanolamines 18 and 19: DGcs, NCC bond angles, OH bond lengths and NPA charges on the N atoms. Compound Entry
DGcs [kcal mol1]
ff NCC bond angle [8]
r(OH) []
NPA charge on N atom
monoethanolamine (MEA) (1) 3-amino-1-propanol (4) 2-piperidinemethanol (18) 3-piperidineethanol (19)
5.83 4.96 4.67 3.00
113.4 – 114.1 –
1.630 1.640 1.685 1.699
0.882 0.877 0.738 0.729
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CHEMPHYSCHEM ARTICLES the steric hindrance around the carbamate nitrogen, the stability goes down and thereby the free amine concentration in the solution increases, which leads to an increase in the CO2 absorption capacity of the alkanolamine.
www.chemphyschem.org potential candidates as industrial solvents for CO2 capture and could be high-capacity replacements for the currently most frequently used solvent, monoethanolamine (MEA).
Acknowledgements 2.5. Advantages of Sterically Hindered Amines for CO2 Capture From Sections 2.1–2.4, it follows that some of the sterically hindered amines may display advantages for CO2 capture. The ideal solvent for CO2 capture should exhibit a high boiling point, form relatively unstable carbamates with CO2, display low regeneration costs and, most importantly, high CO2 absorption rates. The main advantage of sterically hindered amines (Table 5) is that they form less stable carbamates than MEA. A reduction in the carbamate stability leads to a higher free amine concentration in solution, thereby increases the CO2 absorption capacity of the amine. Therefore, high CO2 absorption capacities are expected with secondary alkanolamines with sterically hindering groups near the N atom, like 2-(1methylethylamino)ethanol (8), 2-(2-methyl)propylamino)ethanol (9), 2-N-methylamino-2-methyl-1-propanol (11) and 2-Nethylamino-2-methyl-1-propanol (12). The boiling points of compounds 1, 8, 9, 11 and 12 were estimated to be 170, 183, 175, 199 and 193 8C, respectively. The higher boiling points of compounds 8, 9, 11 and 12 as compared to 1 indicate that the use of these compounds for CO2 capture can decrease volatility losses, which are important for regeneration of the amine. However, although 2-amino-2-methyl-1,3-propanediol (16) forms a less stable carbamate than MEA, it has the disadvantage that it has a lower boiling point than MEA. Based on the carbamate stability towards hydrolysis and the boiling points, secondary alkanolamines with sterically hindering groups near the nitrogen atoms could be highly promising solvents for CO2 capture.
This research is part of the CATO-2 programme, the Dutch national research programme on CO2 capture and storage, funded by the Dutch Ministry of Economic Affairs. We thank the collaborating TNO (Dr. Earl Goetheer) and TU Delft teams (Prof. Thijs Vlugt and Sayee Prasaad Balaji) for stimulating discussions. Keywords: alkanolamines · carbamate stability computational chemistry · CO2 capture · cyclic amines
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[1] IPCC “Summary for Policymakers” Climate Change 2007. [2] a) D. P. Schrag, Science 2007, 315, 812 – 813; b) E. Kintisch, Science 2007, 317, 184 – 186. [3] A. B. Rao, E. S. Rubin, Environ. Sci. Technol. 2002, 36, 4467 – 4475. [4] a) R. Idem, M. Wilson, P. Tontiwachwuthikul, A. Chakma, A. Veawab, A. Aroonwilas, D. Gelowitz, Ind. Eng. Chem. Res. 2006, 45, 2414 – 2420; b) N. McCann, M. Maeder, M. Attalla, Ind. Eng. Chem. Res. 2008, 47, 2002 – 2009; c) G. Sartori, D. W. Savage, Ind. Eng. Chem. Fundam. 1983, 22, 239 – 249. [5] G. T. Rochelle, Science 2009, 325, 1652 – 1654. [6] a) A. Henni, J. Li, P. Tontiwachwuthikul, Ind. Eng. Chem. Res. 2008, 47, 2213 – 2220; b) M. Kundu, S. S. Bandyopadhyay, Fluid Phase Equilib. 2006, 248, 158 – 167; c) S. Ma’mun, V. Y. Dindore, H. F. Svendsen, Ind. Eng. Chem. Res. 2007, 46, 7849 – 7850; d) B. P. Mandal, S. S. Bandyopadhyay, Chem. Eng. Sci. 2006, 61, 5440 – 5447; e) N. Ramachandran, A. Aboudheir, R. Idem, P. Tontiwachwuthikul, Ind. Eng. Chem. Res. 2006, 45, 2608 – 2616; f) A. K. Saha, S. S. Bandyopadhyay, A. K. Biswas, Chem. Eng. Sci. 1995, 50, 3587 – 3598; g) C. S. Ume, E. Alper, Turk. J. Chem. 2012, 36, 427 – 435; h) S. Xu, Y. W. Wang, F. D. Otto, A. E. Mather, Chem. Eng. Sci. 1996, 51, 841 – 850; i) X. Zhang, C. F. Zhang, Y. Liu, Ind. Eng. Chem. Res. 2002, 41, 1135 – 1141. [7] T. Schwabe, S. Grimme, Acc. Chem. Res. 2008, 41, 569 – 579. [8] R. Peverati, D. G. Truhlar, J. Phys. Chem. Lett. 2012, 3, 117 – 124. [9] a) A. V. Marenich, C. J. Cramer, D. G. Truhlar, J. Phys. Chem. B 2009, 113, 6378 – 6396; b) A. V. Marenich, R. M. Olson, C. P. Kelly, C. J. Cramer, D. G. Truhlar, J. Chem. Theory Comput. 2007, 3, 2011 – 2033. [10] E. F. da Silva, H. F. Svendsen, Ind. Eng. Chem. Res. 2006, 45, 2497 – 2504. 3. Conclusions [11] T. A. Halgren, J. Comput. Chem. 1996, 17, 520 – 552. [12] H. Black, Chem. World 2008, 5, 64 – 65. Quantum chemical calculations indicate that important param[13] Gaussian 09, Revision D.01, M. J. Frisch, G. W. Trucks, H. B. Schlegel, G. E. eters that influence the carbamate stability of the alkanolaScuseria, M. A. Robb, J. R. Cheeseman, G. Scalmani, V. Barone, B. Mennucci, G. A. Petersson, H. Nakatsuji, M. Caricato, X. Li, H. P. Hratchian, mines are: steric hindrance around the N atom, charge on the A. F. Izmaylov, J. Bloino, G. Zheng, J. L. Sonnenberg, M. Hada, M. Ehara, N atom and intramolecular hydrogen bonding. Secondary alkaK. Toyota, R. Fukuda, J. Hasegawa, M. Ishida, T. Nakajima, Y. Honda, O. nolamines with moderately sterically hindering groups near Kitao, H. Nakai, T. Vreven, J. A. Montgomery, Jr., J. E. Peralta, F. Ogliaro, the N atoms form less stable carbamates, and are thus highM. Bearpark, J. J. Heyd, E. Brothers, K. N. Kudin, V. N. Staroverov, R. Kobayashi, J. Normand, K. Raghavachari, A. Rendell, J. C. Burant, S. S. Iyengar, J. Tomasi, M. Cossi, N. Table 5. Sterically hindered amines with their boiling points, pKa and DGcs values. Rega, J. M. Millam, M. Klene, J. E. Knox, J. B. Cross, V. Bakken, C. Compound Boiling Point [8C] pKa DGcs[kcal mol1] Adamo, J. Jaramillo, R. Gomperts, [25] monoethanolamine (MEA) (1) 170 9.51 5.83 R. E. Stratmann, O. Yazyev, A. J. 1.56 2-(1-methylethylamino)ethanol (8) 183 10.43[b] Austin, R. Cammi, C. Pomelli, J. W. 10.32[b] 2.85 2-(2-methyl)propylamino)ethanol (9) 199a Ochterski, R. L. Martin, K. Moroku10.14[b] 1.41 2-N-methylamino-2-methyl-1-propanol (11) 175[a] ma, V. G. Zakrzewski, G. A. Voth, P. 9.97[b] 0.84 2-N-ethylamino-2-methyl-1-propanol (12) 193[a] Salvador, J. J. Dannenberg, S. Dap[25] 1.90 2-amino-2-methyl-1,3-propanediol (16) 151 8.84 prich, A. D. Daniels, . Farkas, J. B. Foresman, J. V. Ortiz, J. Cioslowski, [a] Boiling points are calculated with ACD labs software. [b] pKa values are calculated with SMD/SCS-MP2/6D. J. Fox, Gaussian, Inc., Walling311 + G(2d,2p)//SCS-MP2/6-311 + G(2d,2p) calculations. ford CT, 2009.
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CHEMPHYSCHEM ARTICLES [14] Y. Shao, L. F. Molnar, Y. Jung, J. Kussmann, C. Ochsenfeld, S. T. Brown, A. T. B. Gilbert, L. V. Slipchenko, S. V. Levchenko, D. P. O’Neill, R. A. DiStasio, Jr., R. C. Lochan, T. Wang, G. J. O. Beran, N. A. Besley, J. M. Herbert, C. Y. Lin, T. Van Voorhis, S. H. Chien, A. Sodt, R. P. Steele, V. A. Rassolov, P. E. Maslen, P. P. Korambath, R. D. Adamson, B. Austin, J. Baker, E. F. C. Byrd, H. Dachsel, R. J. Doerksen, A. Dreuw, B. D. Dunietz, A. D. Dutoi, T. R. Furlani, S. R. Gwaltney, A. Heyden, S. Hirata, C.-P. Hsu, G. Kedziora, R. Z. Khalliulin, P. Klunzinger, A. M. Lee, M. S. Lee, W. Liang, I. Lotan, N. Nair, B. Peters, E. I. Proynov, P. A. Pieniazek, Y. M. Rhee, J. Ritchie, E. Rosta, C. D. Sherrill, A. C. Simmonett, J. E. Subotnik, H. L. Woodcock III, W. Zhang, A. T. Bell, A. K. Chakraborty, D. M. Chipman, F. J. Keil, A. Warshel, W. J. Hehre, H. F. Schaefer III, J. Kong, A. I. Krylov, P. M. W. Gill, M. Head-Gordon, Phys. Chem. Chem. Phys. 2006, 8, 3172 – 3191. [15] R. Ditchfield, W. J. Hehre, J. A. Pople, J. Chem. Phys. 1971, 54, 720 – 723. [16] a) C. Møller, M. S. Plesset, Phys. Rev. 1934, 46, 618; b) A. E. Reed, R. B. Weinstock, F. Weinhold, J. Chem. Phys. 1985, 83, 735 – 746. [17] a) S. Gangarapu, A. T. M. Marcelis, H. Zuilhof, ChemPhysChem 2012, 13, 3973 – 3980; b) S. Gangarapu, A. T. M. Marcelis, H. Zuilhof, ChemPhysChem 2013, 14, 990 – 995.
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www.chemphyschem.org [18] E. F. da Silva, H. F. Svendsen, Ind. Eng. Chem. Res. 2003, 42, 4414 – 4421. [19] S. M. Bachrach, J. Org. Chem. 2008, 73, 2466 – 2468. [20] F. H. Allen, O. Kennard, D. G. Watson, L. Brammer, A. G. Orpen, R. Taylor, J. Chem. Soc. Perkin Trans. 2 1987, S1 – S19. [21] N. McCann, M. Maeder, H. Hasse, J. Chem. Thermodyn. 2011, 43, 664 – 669. [22] I. V. Alabugin, M. Manoharan, S. Peabody, F. Weinhold, J. Am. Chem. Soc. 2003, 125, 5973 – 5987. [23] S. Paul, A. K. Ghoshal, B. Mandal, Ind. Eng. Chem. Res. 2009, 48, 1414 – 1419. [24] P. Singh, J. P. M. Niederer, G. F. Versteeg, Int. J. Greenhouse Gas Control 2007, 1, 5 – 10. [25] E. P. L. Hunter, S. G. Lias, J. Phys. Chem. Ref. Data 1998, 27, 413 – 656.
Received: September 4, 2013 Revised: October 5, 2013 Published online on November 6, 2013
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Carbamate Stabilities of Sterically Hindered Amines from Quantum Chemical Methods: Relevance for CO2 Capture Satesh Gangarapu,[a] Antonius T. M. Marcelis,[a] and Han Zuilhof*[a, b] The influence of electronic and steric effects on the stabilities of carbamates formed from the reaction of CO2 with a wide range of alkanolamines was investigated by quantum chemical methods. For the calculations, B3LYP, M11-L, MP2, and spincomponent-scaled MP2 (SCS-MP2) methods were used, coupled with SMD and SM8 solvation models. A reduction in carbamate stability leads to an increased CO2 absorption capacity of the amine and a reduction of the energy required for sol-
vent regeneration. Important factors for the reduction of the carbamate stability were an increase in steric hindrance around the nitrogen atom, charge on the N atom and intramolecular hydrogen bond strength. The present study indicates that secondary ethanolamines with sterically hindering groups near the N atom show significant potential as candidates for industrial CO2-capture solvents.
1. Introduction The increased emission of anthropogenic greenhouse gases (GHGs) has the potential to significantly increase the average global temperature.[1] Carbon dioxide is considered one of the major contributors to global warming, due to its abundance compared to other greenhouse gases.[2] Therefore, reduction of the atmospheric emission of CO2 is considered a primary target in environmental policies. Absorption of CO2 with alkanolamines in mixtures with absorption promoters has traditionally been used for its removal from exhaust gases. Among the available alkanolamines, monoethanolamine (MEA) is the most commonly used industrial solvent for CO2 capture.[3] This is because MEA has a high reactivity towards CO2, low solvent cost, low molecular weight and good water solubility.[4] However, MEA has certain disadvantages, such as the formation of rather stable carbamates with CO2 and a concomitantly high energy demand for regeneration, a relatively high volatility, and it shows significant long-term thermal and oxidative degradation. These factors make MEA a suboptimal solvent for CO2 capture from power plants.[5] Clearly, there is considerable scope for improving CO2-capturing solvents, which have higher CO2 absorption rates, lower regeneration costs and lower degradation rates than MEA. Recently, sterically hindered primary amines and secondary amines have received much attention as CO2-capturing sol[a] S. Gangarapu, Dr. A. T. M. Marcelis, Prof. H. Zuilhof Laboratory of Organic Chemistry Wageningen University Dreijenplein 8 6703 HB Wageningen (The Netherlands) Tel: (+ 31) 317-482361 E-mail: [email protected] [b] Prof. H. Zuilhof Department of Chemical and Materials Engineering King Abdulaziz University, Jeddah (Saudi Arabia) Supporting Information for this article is available on the WWW under http://dx.doi.org/10.1002/cphc.201300820.
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vents due to their fast reaction rates with CO2 and lower regeneration costs than conventional primary amines like MEA.[6] Such sterically hindered amines form less stable carbamates upon reaction of CO2 and thereby increase the CO2 absorption capacity from 0.5 mole of CO2 per mole of amine to up to 1 mole of CO2 per mole of amine (Scheme 1). Widely used ster-
Scheme 1. Equations describing the possible reactions of alkyl amines and carbamates with CO2 and water.
ically hindered amines for CO2 capture are 2-amino-2-methyl-1propanol (AMP) and diethanolamine (DEA). Despite the advantages provided by their bulkier substituents, these solvents are not ideal due to a poor reactivity towards CO2 and improvements are still needed. In the search for new solvents for CO2 capture and optimal use thereof, we therefore need an improved and more detailed understanding of the reaction of amines with CO2, including an accurate estimate of the factors that determine their corresponding carbamate stability. In the present study, we thus investigate the electronic and steric effects on the carbamate stabilities of four series of alkanolamines, and we do this by calculating the Gibbs free energy for Reaction (C) with quantum chemical methods. The first series are primary alkanolamines in which one or two CH3 groups are present at the a or b position from the amine group (Figure 1: entries 1–5). The second series consists of secondary alkanolamines that have an increasingly bulky alkyl ChemPhysChem 2013, 14, 3936 – 3943
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www.chemphyschem.org and [Eq. (3)]: DDGs ¼ DGs ðRNH2 Þ þ DGs ðHCO3 ÞDGs ðRNHCO2 ÞDGs ðH2 OÞ
Figure 1. Alkanolamines for CO2 capture as investigated in this study. Different series of alkanolamines are separated by dotted lines.
group on the N atom (Figure 1: entries 6–12). To the best of our knowledge, the effects of bulky groups like isopropyl, isobutyl or secondary butyl substituents on the N atom of MEA (1), and N-methyl or N-ethyl substitution on AMP (5) for application in CO2 capture were not studied yet. The third series are secondary alkanolamines with various N-bound substituents like hydroxyalkyl or aminoalkyl with variable chain lengths are present (Figure 1: entries 13–17). Finally, the fourth series consists of the cyclic amines 2-piperidinemethanol and 2-piperidineethanol (Figure 1: entries 18, 19), which have already shown some promise in CO2 capture.[4c] To this aim, B3LYP/6-311 + G(2d,2p), MP2/6-311 + G(2d,2p), spin-component-scaled MP2/6-311 + G(2d,2p) (SCS-MP2)[7] and M11 L/6-31 + G(d,p)[8] calculations were performed to obtain the Gibbs free energy for the carbamate hydrolysis reaction. Solvents effects were mimicked with the universal solvation models SM8 and SMD.[9] Through this approach, we aim to delineate the electronic and steric effects on the carbamate stabilities of sterically hindered amines. Based on these finding we propose some interesting alkanolamines for industrial use as CO2 capture agents.
Computational Methods and Procedures The carbamate stabilities towards hydrolysis were calculated by using the thermodynamic cycle reported by da Silva et al.[10] The reaction energy in solution (DGcs) can be written as Equation (1): DGcs ¼ DGcg þ DDGs where [Eq (2)]: DGcg ¼ Gg ðRNH2 Þ þ Gg ðHCO3 ÞGg ðRNHCO2 ÞGg ðH2 OÞ
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ð1Þ
Gas-phase conformational analysis was done on the species that are described in the thermodynamic cycle to find the lowestenergy conformers. A conformer search was performed with the MMFF94[11] force field in ChemBio3D.[12] The lowest-energy conformers in the solution phase were found to be the same as in the gas phase. The pre-optimized gas-phase geometries obtained with MMFF94 were further optimized in solution phase with B3LYP/6311 + G(2d,2p), MP2/6-311 + G(2d,2p) and M11-L/6-31 + G(d,p) calculations. Frequency analyses were carried out on the optimized geometries with the same basis set to find true minima on the potential energy surface. Thermal corrections and entropic contributions were added to the total electronic energy for the calculation of the free energy. The universal solvation models (SMD and SM8) that we used for the calculation of the free energy of solvation are trained against the experimental free energies of solvation, and include an entropic contribution to the free energy of solvation. All calculations were carried out with the Gaussian 09[13] software package, Q-chem 4.0[14] and Gaussview 5.0 visualization software. B3LYP, MP2, SCS-MP2 calculations were carried out using Gaussian 09 with the use of the SMD solvent model with the 6-311 + G(2d,2p) basis set.[15] In case of M11L, calculations were carried out using Q-Chem 4.0 [which allows the use of the SM8 solvent model with basis sets up to 6-31 + G(d,p)]. Recently, it was shown that SCS-MP2 methods give significant improvements in cases where MP2 underperforms, which prompted us to use SCS-MP2 energies wherever possible.[7, 16] All data in the tables below represent SMD/ B3LYP/6-311 + G(2d,2p)//SMD/B3LYP/6-311 + G(2d,2p), SMD/SCSMP2/6-311 + G(2d,2p)//SMD/MP2/6-311 + G(2d,2p) calculations. Data of analogous SM8/M11-L/6-31 + G(d,p)//SM8/M11-L/6-31 + G(d,p) calculations are given in the Supporting Information. Atomic charges of all species were calculated by natural population analysis (NPA)[16b] and hybridizations by NBO analysis (NBO)[16b] on the solution phase-optimized geometries.
2. Results and Discussion The geometries of the alkanolamines and their carbamates were optimized with SCS-MP2/6-311 + G(2d,2p) and M11-L/631 + G(d,p) methods. From previous studies, it was found that both methods yield highly accurate pKa and carbamate stability data and thus display the same trends.[17] The lowest-energy conformers of carbamate anions of selected alkanolamines after MP2-optimization, along with their intramolecular hydrogen bond lengths and NC bond lengths are shown in Figure 3. All other MP2-optimized geometries, together with the M11-optimized geometries of all compounds under study, are shown in the Supporting Information (Figures S1 and S2, respectively). The lowest-energy conformers of the carbamate anions show H-bonds between an O atom of the COO group and the H atom of the OH group. The lowest-energy conformers of the alkanolamines show intramolecular H-bonding between the amine N atom and the H atom of the hydroxyl group. The lowest-energy conformers of either alkanolamine or its carbamate anion in the gas phase and solution phase are almost identical in geometry, and the same intramolecular Hbonds are observed in the carbamate anions in the gas phase and in solution.[18] The Gibbs free energy as calculated from ChemPhysChem 2013, 14, 3936 – 3943
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Figure 2. Thermodynamic cycle used for the calculation of the Gibbs free energy of Reaction (C).
the thermodynamic cycle (Figure 2) was used as a parameter to interpret the carbamate stabilities of the sterically hindered amines.
2.1. Substitution of C-linked Methyl Groups on MEA For the sterically hindered primary amines, the effect of substitution of one or two methyl groups either on the a- or bcarbon atom from the amine functional group was investigated. The lowest-energy conformers of the carbamate anions of relatively important alkanolamines are shown in Figure 3, to-
Figure 3. Lowest-energy conformers of the carbamate anions formed upon reaction of CO2 with some relatively important substituted alkanolamines, showing also the O···HO and NC bond lengths [SCS-MP2/6-311 + G(2d,2p) data]. Black: N; dark grey: O; grey: C; white: H.
gether with their H-bond and NC bond lengths. The carbamate stability of aCH3 substituted MEA (3) was found to be lower than the carbamate stability of MEA (1), due to the increase in steric hindrance around the amine group (Table 1). Further substitution
with a second methyl group on the a-position to the amine group yields the well-known alkanolamine AMP (5), which displays a further decreased carbamate stability of 3.59 kcal mol1. The reduction in the intramolecular hydrogen bond length between the hydrogen atom from the hydroxyl group and the oxygen atom from the COO group of a carbamate with one or more CH3 groups on the a-carbon atom is reflected by a decrease in NCC(OH) bond angle, which decreases stepwise from 113.48 via 111.98 to 111.38 in going from MEA-carbamate (1) via the carbamate of 3 to that of 5. However, the N CC(OH) bond angle of 2 is 113.78, which is similar to that of the MEA carbamate (1), indicating that there is almost no effect of the bond angle due to substitution of a CH3 group on the b-carbon atom from the amine group. The decrease in the NCC(OH) bond angle upon substitution of CH3 groups on the a-carbon atom is due to the so-called Thorpe–Ingold effect.[19] A result of a smaller NCC(OH) angle is that the intramolecular (OC)O···H(O) bond becomes shorter. The increase in steric hindrance thus corresponds with a decrease in the intramolecular H-bond length, with (OC)O···H(O) distances for the carbamate of 1 of 1.630 to 1.612 for 3 and 1.589 for 5. The effect of intramolecular hydrogen bonding in the a-CH3 substituted alkanolamines is further studied by looking at the structural changes between the carbamate of 5 with and without intramolecular hydrogen bonding. The NCC(OH) bond angle of 5 with hydrogen bonding was found to be 111.38, which is higher than the NCC(OH) bond angle of 106.78 in the non H-bonded carbamate. The increase in the NCC(OH) bond angle of 5 with hydrogen bonding is accompanied by a decrease in CH3CCH3 bond angle from 109.38 to 107.98. The decrease in the carbamate stability of a-CH3 substituted alkanolamines can be further explained by looking at the NC bond lengths in the carbamate species of substituted MEA. The NC bond length in the carbamate of unsubstituted MEA is 1.443 , which is characteristic for an NC single bond.[20] Upon substitution of a -CH3 groups on MEA, the NC bond length increases (data shown in supporting information Table S2), in line with a weakening of the NC bond by substitution of hydrogen by a -CH3 group. This weakening is crucial for a more cost-effective regenerative desorption of CO2. Both SMD/SCS-MP2 and SM8/M11-L calculations show that DG for the carbamate hydrolysis reaction becomes less negative with an increase in steric hindrance around the N atom. The carbamate stability towards hydrolysis of substituted alkanolamines follows the order: MEA > 1-amino-2-propanol > 2amino-1-propanol > 3-amino-1-propanol > AMP, which follows
Table 1. Alkyl-substituted MEAs: DGcs, NCC bond angle, OH bond lengths and NPA charges on the N atoms. Compound Entry
DGcs [kcal mol1]
ffNCC [8]
r(OH) []
NPA charge on N atom
monoethanolamine (MEA) (1) 1-amino-2-propanol (2) 2-amino-1-propanol (3) 3-amino-1-propanol (4) 2-amino-2-methyl-1-propanol (AMP) (5)
5.83 5.34 5.04 4.96 3.59
113.4 113.7 111.9 – 111.3
1.630 1.618 1.612 1.640 1.589
0.882 0.879 0.879 0.877 0.865
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the same trend as experimentally found for compounds for which these data are available.[4c] The calculated DGcs of MEA carbamate was 5.83 kcal mol1, 6.79 kcal mol1, 6.28 kcal mol1 and 6.55 kcal mol1 with SCS-MP2, B3LYP, MP2 and M11-L respectively, that is, in good agreement with the experimentally reported value of 4.92 kcal mol1 for the reverse reaction.[21] Although accurate experimental data are scarce, this suggests that the methods used (SMD/SCS-MP2, SMD/MP2, SMD/B3LYP and SM8/M11-L) can accurately calculate the Gibbs free energy for the carbamate hydrolysis reaction. To check for dependence of the overall result on the solvent model and basis set, we performed the calculation on three compounds from different series that is, 1, 8, and 18, with SMD/M11-L/6-311 + G(2d,2p) (data shown in Table S1) and compared it with one used in the calculations that is, SM8/M11-L/6-31 + G(d,p). It was found that the results from both models are almost identical. The DGcs values from B3LYP, M11 L and MP2 deviate from the SCS-MP2 values by ~ 1.2 kcal mol1, ~ 1.0 kcal mol1 and ~ 0.5 kcal mol1, respectively. However, the results from SCSMP2 are more accurate than those from M11-L, and therefore the results from SCS-MP2 are discussed here and data from B3LYP, MP2 and M11-L are shown in Tables S1 and S2. A good correlation (R2 = 0.96) was observed between DGcs for the carbamate hydrolysis reaction and the charge on the N atom in alkanolamines (Figure 4 a). It clearly indicates that the higher the charge on the N atom, the stronger the bond to CO2. A very good correlation (R2 = 0.99) was observed between DGcs and the O···HO intramolecular hydrogen bond length in carbamates (Figure 4 b). 2.2. Substitution of N-Linked Alkyl Groups on MEA In this section, the direct substitution of an H atom on the amine N atom in MEA by a variety of increasingly bulky alkyl groups is investigated. The DGcs, intramolecular H-bond lengths, NCC bond angles and NPA charge on the N atom of the carbamate anions of N-alkyl-substituted MEAs are shown in Table 2. As expected, direct substitution of a CH3 group on the N atom of the amine functional group (6) decreases the carbamate stability of the alkanolamine, as seen from the decrease in the Gibbs free energy of reaction from 5.83 kcal mol1 to
Figure 4. Correlation plots between DGcs for the carbamate hydrolysis reaction and a) the NPA charge on the N atom, or b) DGcs the O···HO intramolecular H-bond length in the carbamates of analogous compounds (1–3 and 5) (4 is excluded from the graph, because it is an aminopropanol).
4.80 kcal mol1 (Table 2). Substitution by an ethyl group (7) on the N atom yielded a DGcs of 4.53 kcal mol1 (Table 2), which is lower than for the alkanolamines 1 and 6. This is obviously due to the inductive effect of a CH3 group on the N atom. The total NPA charge of the CH3 group in the carbamate of 6 is 0.357, which is more positive than the NPA charge on the hydrogen atom (0.206) in carbamate 1. Further increase in steric hindrance around the N atom by substitution with an isopropyl group yielded the desired clear decrease in carbamate stability. The DGcs of the isopropyl group substituted alkanolamine 8 was found to be ~ 4.3 kcal mol1 smaller (i.e., less negative) than of unsubstituted MEA (1): only 1.56 kcal mol1 (Table 2). Surprisingly, substitution by more bulky groups on the N atom, like isobutyl (9) or secondary butyl (10), resulted in a more negative Gibbs free energy than obtained for the isopropyl group.
Table 2. N-substituted MEAs and 2-aminopropanols: DGcs, NCC bond angle, OH bond lengths and NPA charges on the N atoms. Compound Entry
DGcs [kcal mol1]
ffNCC [8]
r(OH) []
NPA charge on N atom
2-(methylamino)ethanol (6) 2-(ethylamino)ethanol (7) 2-(1-methyl)ethylamino) ethanol (8) 2-(2-methyl)propylamino) ethanol (9) 2-(1-methylpropylamino) ethanol (10) 2-(N-methylamino)-2-methyl-1-propanol (11) 2-(N-ethylamino)-2-methyl1-propanol (12)
4.80 4.53 1.56
113.6 113.9 114.1
1.622 1.630 1.646
0.738 0.736 0.729
2.85
114.0
1.642
0.731
3.53
114.0
1.638
0.735
1.41 0.84
112.2 111.6
1.575 1.568
0.744 0.740
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CHEMPHYSCHEM ARTICLES We hypothesized that this might be due to the difficulty of solvent accessibility during the solvent energy calculations with the bulky groups. To test this idea, the solvation energies for alkanolamines 6–10 were calculated with inclusion of one explicit water molecule in the solvation energy calculations. The results of including one explicit water molecule are shown in Figure 5. The distance between the N atom of the carba-
Figure 5. Optimized structures obtained by inclusion of an explicit water molecule to the reaction center (N atom) in alkanolamines 6–10. Black: N; dark grey: O; grey: C; white: H.
mate and O atom of this solvating water molecule is 2.334 for 6, 2.626 for 7 and 3.192 for 8. For alkanolamine 10, which contains the bulky secondary butyl group, the distance is even > 4 . This clearly indicates that upon increasing the steric hindrance around the nitrogen atom in alkanolamines, it becomes more difficult for the solvent to interact with the reaction center. This observation is further supported by calculating the solvation energies of carbamates of 6–10. The solvation energies for secondary butyl (10) and isobutyl (9) substituted MEA are ca. 2–3 kcal mol1 smaller than the solvation energy for the sterically less hindered compounds 6–8 (full data shown in Table S4). Therefore, the steric hindrance around the N atom seems to have an optimum when aiming for an improved CO2 absorption capacity, that is, with the isopropyl group. By comparing DGcs of the carbamates of ethanolamines with different substituents on the N atom, it is seen that substitution of a H atom by either a methyl, ethyl or isopropyl group improves the CO2 absorption capacity by a decrease in the stability of the carbamate anions. The changes in the intramolecular H-bond length in N-alkyl MEA nicely correspond with the changes in NCC angles. Substitution of an ethyl group (7) on the N atom increased ffNCC from 113.48 for MEA to 113.98 for 7, and increased the intramolecular H-bond length from 1.620 to 1.630 . Further substitution of an isopropyl group on MEA (8) increased ffNCC to 114.18 and the intramolecular H-bond length to 1.646 .
2013 Wiley-VCH Verlag GmbH & Co. KGaA, Weinheim
www.chemphyschem.org A good correlation (R2 = 0.94) was observed between DGcs and the charge on the N atom in the ethanolamines (Figure 6 a). This also clearly indicates that the higher the charge
Figure 6. Correlation plots between a) the Gibbs free energy for the carbamate hydrolysis reaction (DGcs) and NPA charge on N atom, and b) DGcs and O···HO intramolecular hydrogen bond length in carbamates of ethanolamines 6–10.
on the N atom, the stronger its binding to CO2. A reasonably good correlation (R2 = 0.89) was observed between DGcs and the O···HO intramolecular H-bond length in carbamates (Figure 6 b). It was found that a longer H-bond length, that is, weakened H-bond in the reactant, decreases the resulting carbamate stability, which is in contrast with the observation mentioned in Section 2.2. This difference is due to the direct substitution of alkyl groups on the N atom in these molecules. Due to the increasing steric hindrance, the bond angles around the N atoms increase to close to 1208 in compounds 9 and 10, corresponding to a hybridization change from sp3 to almost sp2.[22] The change in the hybridization around the N atom was further confirmed by NBO analysis of the carbamates of 9 and 10. The correlation plot between DGcs and the charge on the nitrogen atom in all alkanolamines included in this study is shown in Figure S3, which shows that DGcs and the NPA charge depend strongly on the nature of the amine group in alkanolamines. Methyl substitution on the N atom of AMP (11) decreased the Gibbs free energy of reaction from 3.59 kcal mol1 to 1.41 kcal mol1 (Table 2). Furthermore, substitution of an ethyl group on AMP (12) decreased the Gibbs free energy of reaction to 0.84 kcal mol1, also due to poor solvation (Table 2). This clearly indicates that upon increase in the steric hindrance around the N in the alkanolamine, the carbamate stability decreases, thereby increasing the CO2 absorption capacity of the ChemPhysChem 2013, 14, 3936 – 3943
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alkanolamine. This means that substitution of one of the hydrogen atoms on the N of AMP by either a methyl or ethyl group provides a highly promising route to reduce the energy needed for regeneration of alkanolamines from carbamates.
2.3. Substitution of Hydroxyalkyl and Aminoalkyl Functional Groups on the N Atom in MEA In this section, the carbamate stabilities of a series of alkanolamines that contain alkylhydroxy or alkylamine functional groups (13–17) were calculated and compared with MEA (1). The DGcs, intramolecular H-bond lengths, ffNCC and NPA charge on the N atom of carbamate anions of all alkanolamines (13–17) are shown in Table 3. Upon substitution of an N-bound H atom by an N-hydroxyethyl (CH2CH2OH) group (13) or an aminoethyl (CH2 CH2NH2) group (14), DGcs decreased from 5.83 kcal mol1 to 4.76 and 4.51 kcal mol1, respectively. This decrease in carbamate stability is attributed to the electron-withdrawing nature of these groups. The formation of less stable carbamates from 13 and 14 is further supported by the NPA charge on the nitrogen atom. The NPA charges on the N atoms in both 13 (0.749) and 14 (0.734), are smaller than the value of 0.882 found for MEA (1), indicating a reduced reactivity towards CO2. Compounds 14 and 17 contain two reaction centers that is, both primary and secondary amine groups. We considered the formation of a carbamate on the secondary amine group, due its higher reactivity towards CO2 than a primary amine group. In compound 15, in which two 2-hydroxypropyl moieties are present on the N atom, the carbamate stability decreases due to the increase in the steric hindrance. Compound 16 forms the least stable carbamate among all amines in Table 3, despite having a primary amine group. This is caused by the formation of a second H-bond (OH···N) in this compound. This is further indicated by the longest NC bond length (data shown in the Supporting Information) in the carbamate anion of compound
16. A very poor correlation was observed between DGcs either with the NPA charge on the N atom or with the intramolecular H-bond length. This may be due to the presence of extra intramolecular H-bonding in the alkanolamines 13–17. Data from both calculations (SCS-MP2 and M11-L) show that the carbamate stabilities decrease in the following order: MEA > DEA > AMP, which is the same as the experimental values reported by Satori and Savage.[4c] It clearly confirms that both methods (SMD/SCS-MP2 and SM8/M11-L) are good enough to calculate the carbamate stabilities of alkanolamines accurately.
2.4. Linear Alkanolamines versus Cyclic Alkanolamines In this section, the carbamate stabilities of linear and cyclic alkanolamines are investigated. Cyclic alkanolamines were found to exhibit higher CO2 absorption capacities than conventional alkanolamines.[4c, 23] Table 4 shows the DGcs, H-bond lengths, NCC bond angles and the NPA charge on the N atom of the carbamate anions of cyclic alkanolamines 18 and 19 and of the corresponding linear forms 1 and 4. It was found that DGcs of the carbamates of the linear alkanolamines are higher than of the cyclic alkanolamines with the same number of carbon atoms between the hydroxyl and amine groups. Upon increasing the carbon chain length between the amine and hydroxyl functional groups, the carbamate stability decreased in both linear and cyclic alkanolamines. This is caused by a reduction of the electron withdrawing effect of the hydroxyl group on the nitrogen atom of the amine.[24] In conclusion, the carbamate stability of substituted alkanolamines does not only depend on the different electron-donating or -withdrawing properties of the functional groups, but also depends on the position (a or b or direct) of a substituent with respect to the amine or hydroxyl functional group or linearity of the molecule or the carbon chain length between the amine and hydroxyl functional groups in the alkanolamines. From the calculations we can conclude that upon increasing
Table 3. Hydroxyalkyl- and aminoalkyl substituted MEAs: DGcs, NCC bond angle and OH bond lengths, and NPA charges on the N atoms. Compound Entry
DGcs [kcal mol1]
ffNCC [8]
r(OH) []
NPA charge on N atom
di-(2-hydroxyethyl) amine (13) 2-(2-aminoethyl)amino ethanol (14) 1-(2-hydroxypropyl) aminopropanol (15) 2-amino-2-methyl-1,3-propanediol (16) 2-(3-aminopropyl) aminoethanol (17)
4.76
114.1
1.689
0.749
4.51
113.9
1.633
0.734
4.23 1.90 4.49
114.6 114.2 114.1
1.668 1.674 1.643
0.744 0.882 0.729
Table 4. Linear alkanolamines 1 and 4 and cyclic alkanolamines 18 and 19: DGcs, NCC bond angles, OH bond lengths and NPA charges on the N atoms. Compound Entry
DGcs [kcal mol1]
ff NCC bond angle [8]
r(OH) []
NPA charge on N atom
monoethanolamine (MEA) (1) 3-amino-1-propanol (4) 2-piperidinemethanol (18) 3-piperidineethanol (19)
5.83 4.96 4.67 3.00
113.4 – 114.1 –
1.630 1.640 1.685 1.699
0.882 0.877 0.738 0.729
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CHEMPHYSCHEM ARTICLES the steric hindrance around the carbamate nitrogen, the stability goes down and thereby the free amine concentration in the solution increases, which leads to an increase in the CO2 absorption capacity of the alkanolamine.
www.chemphyschem.org potential candidates as industrial solvents for CO2 capture and could be high-capacity replacements for the currently most frequently used solvent, monoethanolamine (MEA).
Acknowledgements 2.5. Advantages of Sterically Hindered Amines for CO2 Capture From Sections 2.1–2.4, it follows that some of the sterically hindered amines may display advantages for CO2 capture. The ideal solvent for CO2 capture should exhibit a high boiling point, form relatively unstable carbamates with CO2, display low regeneration costs and, most importantly, high CO2 absorption rates. The main advantage of sterically hindered amines (Table 5) is that they form less stable carbamates than MEA. A reduction in the carbamate stability leads to a higher free amine concentration in solution, thereby increases the CO2 absorption capacity of the amine. Therefore, high CO2 absorption capacities are expected with secondary alkanolamines with sterically hindering groups near the N atom, like 2-(1methylethylamino)ethanol (8), 2-(2-methyl)propylamino)ethanol (9), 2-N-methylamino-2-methyl-1-propanol (11) and 2-Nethylamino-2-methyl-1-propanol (12). The boiling points of compounds 1, 8, 9, 11 and 12 were estimated to be 170, 183, 175, 199 and 193 8C, respectively. The higher boiling points of compounds 8, 9, 11 and 12 as compared to 1 indicate that the use of these compounds for CO2 capture can decrease volatility losses, which are important for regeneration of the amine. However, although 2-amino-2-methyl-1,3-propanediol (16) forms a less stable carbamate than MEA, it has the disadvantage that it has a lower boiling point than MEA. Based on the carbamate stability towards hydrolysis and the boiling points, secondary alkanolamines with sterically hindering groups near the nitrogen atoms could be highly promising solvents for CO2 capture.
This research is part of the CATO-2 programme, the Dutch national research programme on CO2 capture and storage, funded by the Dutch Ministry of Economic Affairs. We thank the collaborating TNO (Dr. Earl Goetheer) and TU Delft teams (Prof. Thijs Vlugt and Sayee Prasaad Balaji) for stimulating discussions. Keywords: alkanolamines · carbamate stability computational chemistry · CO2 capture · cyclic amines
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Received: September 4, 2013 Revised: October 5, 2013 Published online on November 6, 2013
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