CHEMPHYSCHEM ARTICLES DOI: 10.1002/cphc.201300952
Ab Initio Study of the Adsorption of CO2 on Functionalized Benzenes Maria G. Frysali, Emmanuel Klontzas, and George. E. Froudakis*[a] The interaction of carbon dioxide with a series of functionalized aromatic molecules was studied by using quantum mechanical methods (MP2), to examine the effect of the substituent on the adsorption of CO2. Several different initial configurations of CO2 were taken into account for each functionalized benzene to locate the energetically most favorable configuration. To get a better estimation of the binding energies, we applied an extrapolation scheme to approach the complete basis set. CH2N3-, COOH-, and SO3H-functionalized benzenes were
found to have the strongest interaction with CO2, and the corresponding binding energies were calculated to be 3.62, 3.65, and 4.3 kcal mol1, respectively. Electrostatic potential maps of the functionalized benzenes and electron redistribution density plots of the complexes were also created to get a better insight into the nature of the interaction of CO2 with the functionalized benzenes. The functional groups that were examined can be potentially incorporated in organic bridging molecules that connect the inorganic corners in MOF.
1. Introduction Carbon dioxide is the fourth most abundant gas that is naturally present in the Earth’s atmosphere. Over the past two centuries its concentration in the atmosphere has greatly increased, mainly due to human activities, such as fossil fuel burning. Today, it is considered to be one of the greenhouse gases that contribute to the Earth’s global warming phenomenon. For the next few years, our energy demands will be satisfied by using carbon-based fuels as energy resources. To reduce the amount of CO2 that is released into the atmosphere, the global community is trying to find efficient ways to reduce the amount of CO2 released by using either direct capture from the air or from the resources from which it is produced, such as power plants or combustion engines. The whole process of limiting CO2 emissions due to human activities in the atmosphere is known as the carbon dioxide capture and storage (CCS) procedure.[1] This procedure includes a group of technologies for the capture of CO2 from power plants, followed by compression and transport to permanent storage sites. The capture process comprises the main cost of CCS schemes and represents approximately two thirds of the total cost. There are three processes that are used to capture CO2, known as post-combustion, pre-combustion, and oxyfuel processes. A diverse range of materials with different properties is used for capturing CO2 in these processes. Post-combustion methods have been considered to be less energy-de[a] M. G. Frysali, Dr. E. Klontzas, Prof. G. E. Froudakis Department of Chemistry University of Crete POBOX 2208 71003 Voutes Heraklion (Greece) Fax: (+ 30) 2810545001 E-mail: [email protected] Supporting Information for this article is available on the WWW under http://dx.doi.org/10.1002/cphc.201300952.
2014 Wiley-VCH Verlag GmbH & Co. KGaA, Weinheim
manding methods and a lot of efforts have been devoted to finding suitable materials for this process. Many materials have been investigated for CO2 adsorption. Examples of these materials include physical absorbents,[2] ionic liquids,[3, 4] carbonbased sorbents,[5, 6] metal oxides,[7] hydrotalcite-like[8] compounds, zeolites,[2, 9–15] organic solids,[16, 17] and metal–organic frameworks. In the chemical industry, the separation of CO2 from mixtures with natural gas and flue gas is tremendously important. Natural gas is an alternative substitute for environmentally unfriendly fossil fuels. However, impurities in natural gas, such as CO2, reduce the calorie content and removal of CO2 from natural gas is expensive. For this reason, significant effort has been made in recent years to improve CO2 capture. Although a variety of technologies and methods have been developed, the discovery of new materials with high selectivity toward CO2 in a mixture of gases and high capacities remains a challenge. Of the different types of materials studied for this purpose, metal–organic frameworks (MOFs), which are organic–inorganic hybrid materials composed of metal ions or clusters interconnected through an organic linker, have attracted wide scientific interest owing to their enormous structural and chemical diversity and their potential applications in gas storage, ion exchange, molecular separation, and heterogeneous catalysis.[18–25] Recently, attention has turned to experimental and computational screening studies of different materials to assess CO2 adsorption. A large variety of MOFs, such as IRMOF1, IRMOF3, IRMOF6, IRMOF11, MOF2, MOF74, MOF177, MOF505, and [Cu3(BTC)2], were experimentally examined for their structural and porous attributes towards CO2 adsorption.[26] MOF-177, with a surface area (SA) of 4500 m2 g1, exhibits the highest capacity for CO2 and takes up 33.5 mol kg1 at 32 bar and
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298 K,[27] whereas MOF200 and MOF210[28] achieved CO2 caRegarding previous computational studies,[43–46] density func1 pacities of 54.5 mol kg at 50 bar and 298 K. tional theory (DFT) and post-Hartree–Fock methods[47–51] have At high pressures, CO2 capacities depend on the pore been applied to study many properties of these MOF materials. Traditional DFT functionals fail to evaluate the dispersion intervolume and surface areas[29, 30] of the MOFs. The increase in action, which plays an important role in weakly interacting pore volume and surface area enhance the storage capacity. In complexes, crystal packing of organic molecules, and three-dicontrast, CO2 capacities depend on the heats of adsorption for mensional structures of biological systems. Another disadvantCO2. Many studies have tried to improve the selectivity toward age of this method is the incorrect evaluation of polarization CO2 through polar functional groups, pore-size control,[31–33] and hyperpolarization abilities of large p-conjugated moleopen metal sites,[34–36] or with the introduction of alkali–metal cules.[52] In recent years, DFT and ab initio calculations have cations.[37] been used to study interaction energies and adsorption sites There are studies that indicate that MOFs with highly polar between MOFs and CO2.[53–56] Herein, we used ab initio methligands enhance the storage capacity for CO2. An example is bio-MOF-11, which presents high CO2 uptake at 298 K and very ods to study organic linkers that could improve the CO2 adminor N2 adsorption and is based on the Lewis basic sites of sorption capacity in MOFs. adenine, which includes an amino group and pyrimidine nitrogen atoms with strong affinity for CO2.[31] Another example is 2. Results and Discussion amine-functionalized MIL-53(Al). Amine increases the CO2/CH4 separation factor (from 5 to 60) compared with non-functionalOptimized structures of CO2 with the functionalized benzenes ized MIL-53.[38] Concerning the pore-size effect, there are severcan be seen in Figure 1 and the corresponding binding eneral MOFs with high selectivity for CO2 over N2 or CH4 due to kigies are reported in Table 1. Optimized geometries of the netic separation effect[30, 39] and other MOFs with indicated sestructures that correspond to the energetically most favorable lectivity due to a molecular sieving effect.[29, 40] Open metal structure were located and verified after performing geometry optimizations from different starting geometries. From sites have been also studied for CO2 capture and separation. Figure 1, it can be seen that the optimized structures can be For example, CO2/CH4 selectivities between carborane-based divided into two categories, that is, CO2 can be located above MOFs are compared with and without open metal sites and the benzene ring or it can be located over the functional the results suggested that high selectivity for CO2 over CH4 group. The binding energies calculated at the RI-MP2/def2(17) are obtained.[41] TZVPP level, reported in Table 1, range from 2.27 to In materials that capture carbon dioxide, such as MOFs, the adsorption of CO2 in the pores of the materials plays a major 4.27 kcal mol1 for CO2–C6H5F and CO2–C6H5SO3H, respectively. role. The knowledge of the nature of the interaction of CO2 The largest interaction energy was found when CO2 interactwith the building fragments that form the material is vital in understanding the adsorption process in existing materials and ed with C6H5SO3H, for which the interaction was calculated to in designing new materials for enhanced carbon dioxide capbe 4.27 kcal mol1. In this optimized structure, CO2 is not lo[42] ture. The nature of the interaction of CO2 with molecules has been attributed to the permanent quadrupole moment of CO2. CO2 can act as a Lewis acid or as a Lewis base. It acts as Lewis acid in the presence of electron-rich atoms on substituted benzenes, through interaction with the (electron-deficient) C atom of CO2, or as Lewis base in the presence of electrondeficient atoms of substituted benzenes through interaction with the (electron-rich) O atom of CO2. As has been noted in previous studies,[55, 56] dispersion forces and electrostatic interactions contribute almost equally to the binding energy of CO2 and, in some cases, dispersion forces have a more dominant Figure 1. Optimized structures of C6H5F (1), C6H6 (2), C6H5NO2 (3), C6H5CN (4), C6H5CF3 (5), C6H5OH (6), C6H5PH2 (7), role in the binding energy of the C6H5SH (8), C6H5CH3 (9), C6H5OC2H5 (10), C6H5SO2NH2 (11), C6H5C2N3 (12), C6H5COOH (13), C6H5SO3H (14), and molecular systems. C6H5NH2 (15) with CO2. 2014 Wiley-VCH Verlag GmbH & Co. KGaA, Weinheim
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Table 1. Binding energies (Ecp) and the corresponding O-C-O angles for all CO2 complexes calculated at the RI-MP2/def2-TZVPP level of theory. All energies were calculated by including a counterpoise correction.
C6H5F–CO2 C6H6–CO2 C6H5NO2–CO2 C6H5CN–CO2 C6H5CF3–CO2 C6H5OH–CO2 C6H5PH2–CO2 C6H5SH–CO2 C6H5CH3–CO2 C6H5OC2H5–CO2 C6H5NH2–CO2 C6H5SO2NH2–CO2 C6H5CH2N3–CO2 C6H5COOH–CO2 C6H5SO3H–CO2
Ecp [kcal mol1]
O-C-O angle [8]
2.27 2.34 2.42 2.46 2.55 2.72 2.80 2.92 2.92 3.22 3.34 3.48 3.62 3.65 4.27
179.161 179.024 179.472 178.844 179.418 179.033 179.140 179.132 178.851 178.791 178.484 178.600 178.947 177.988 178.360
cated above the benzene ring, but close to the SO3H functional group. The complex is stabilized by the simultaneous interaction of the carbon atom of CO2 with the oxygen atom of the sulfonate group and the interaction the oxygen of CO2 with the hydrogen atom of the sulfonate group. The CCO2OSO3H and OCO2HSO3H contacts were measured to be 2.90 and 1.98 , respectively. As mentioned above in the Introduction, both the carbon and oxygen atoms of the CO2 can contribute to the stabilization of the complex by acting as a Lewis acid and as a Lewis base, respectively. In this case, the electron-deficient carbon atom of CO2 acts as Lewis acid when interacting with the oxygen atom of the sulfonate group. The oxygen atom of CO2 acts as a Lewis base and forms a hydrogen bond with the electron-deficient hydrogen atom of the sulfonate group, as indicated by the short OCO2HSO3H contact. The next highest binding energy was calculated to be 3.65 kcal mol1 for the interaction of CO2 with C6H5COOH. As in the previous compound, the oxygen of the carboxylic group act as a Lewis base to the electron-deficient carbon atom of the CO2, whereas the hydrogen atom acts as a Lewis acid. The CCO2OCOOH and OCO2HCOOH contacts were measured to be 2.85 and 2.03 , respectively, which are similar to the corresponding values in C6H5SO3H. CO2 is oriented in the same plane, which can be formed from the atoms of the carboxylic group. CO2 interacts with C6H5CH2N3 with a binding energy of 3.62 kcal mol1, which is similar to the binding energy found for the carboxylic group. In this case, CO2 is oriented over CH2N3 in such a way that it can interact with both the nitrogen atom of the azide group and the hydrogen atom of the CH2. Based on the optimized geometry that was obtained from our calculations and the corresponding separations between the atoms of CO2 and the atoms of the CH2N3 group, N3 act as a Lewis base and the formed complex is further stabilized by the formation of a weak hydrogen bond between the oxygen atom of CO2 and the neighboring hydrogen atom of CH2. This is supported by the corresponding separations between the nitrogen atom of N3 with the carbon atom of CO2 (3.09 ) and that of the oxygen atom of the CO2 and the hydrogen atom of CH2 2014 Wiley-VCH Verlag GmbH & Co. KGaA, Weinheim
(2.89 ). The value of 3.09 for the CN contact is a bit longer than the corresponding separation in molecular systems that contain nitrogen atoms that interact with CO2 ( 2.90 ). The binding energy that was obtained for C6H5CH2N3 is larger than the value (2.56 kcal mol1) that was calculated by Vitillo et al.[55] for C6H5N3 with a similar quality of computational methods. This can be attributed to the existence of the weak hydrogen bond that we mentioned previously. The interaction energy of CO2 with C6H5SO2NH2 was calculated to be 3.48 kcal mol1. Again, we have an electron donor– acceptor mechanism for the stabilization of this complex. One of the oxygen atoms of CO2 forms a hydrogen bond with a hydrogen of the NH2 group and the corresponding separation was calculated at 2.18 . The carbon atom of the CO2 acts as a Lewis acid and interacts with the oxygen atom of the SO2NH2 group. The contact between the carbon atom and the oxygen atom is 2.95 , similar to the value found for the CO2– C6H5SO3H complex. The next lowest binding energy was found for the interaction of CO2 with C6H5NH2. The interaction energy for this system was calculated to be 3.34 kcal mol1, in agreement with the value of 3.27 kcal mol1 that was previously found by Vitillo et al.[55] CO2 prefers a location above the N atom of the amino group due to the electron donor–acceptor mechanism between the C atom of CO2 and the N atom of the amino group. We also investigated the interaction of CO2 with C6H5OC2H5. The interaction energy was calculated to be 3.22 kcal mol1. The geometrical configuration of the CO2 is characterized by the formation of two hydrogen bonds between OCO2 and HOC2H5, with a contact of 2.67 for OCO2H and a contact of 2.93 between OCO2 and the H atom of benzene. The C atom of CO2 interacts with the O atom of OC2H5, as can be deduced from their interatomic separation (2.91 ). In the optimized geometries found for CO2 with the other functionalized benzenes, we observed that the CO2 is oriented in a parallel mode above the plane of the benzene ring, with an exception in the case of C6H5CN. In all cases except the latter, CO2 is displaced from the center of mass of the benzene ring, which indicates that there is an additional interaction between the oxygen atoms of the CO2 with the atoms of the polar groups on top of the known interaction between the quadrupole moment of CO2 and the p system of benzene. This phenomenon has also been mentioned in previous studies on the interaction of CO2 with molecular systems that contain aromatic rings. For the interaction of CO2 with C6H5CH3, the calculated value for the interaction was found to be 2.92 kcal mol1. In this case, the methyl group acts as an electron donor to the benzene ring and enhances the interaction of the quadrupole with the p system of the ring. Furthermore, there is an additional contribution due to the interaction between the oxygen of CO2 and the hydrogen from the methyl group, which is reflected in the short contact of 2.66 between these two atoms. The same interaction energy was found also for the interaction of CO2 with C6H5SH. This functional group acts as an electron acceptor and weakens the p system of the ring. The geometry of the optimized structure is stabilized through the interaction of the hydrogen of the SH group with the polarized oxygen of ChemPhysChem 2014, 15, 905 – 911
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a high potential value are represented in red and electron-rich CO2 with an OCO2HSH contact of 2.96 . CO2 interacts with regions that correspond to a low potential are represented C6H5PH2 and C6H5OH with interaction energies of 2.80 and with blue. Electrostatic potential maps can be useful in explain2.67 kcal mol1, respectively. In the case of C6H5OH, CO2 is loing and predicting the geometry of the molecular system and cated on the top of the benzene ring, which means that the the nature of the interactions between two molecules, espemost stable geometry is when CO2 interacts not only with the cially when electrostatic interactions are important in describformation of a weak hydrogen bond between OCO2 and HOH ing their interaction. The stability of the systems that were exbut with the simultaneous interaction of CO2 with the p amined herein comes from the existence of electrostatic intersystem of the ring. The separation between the oxygen atom actions due to the charged nature of CO2 and the functionalof CO2 and the hydrogen atom of the hydroxyl group was measured to be 2.98 . ized benzenes on top of the contribution due to the existence For the other functionalized benzenes that were studied, the of dispersion forces. In fact, previous studies have shown that interaction energies with CO2 ranged from 2.27 to 2.55 kcal dispersion forces and electrostatic interactions contribute almost equally to the binding energy of CO2 and, in some mol1. The interaction energy of CO2 with benzene, for which the location of CO2 is distorted from the center of mass of the cases, dispersion forces have more a dominant role in the binding energy of the molecular systems. ring, was calculated to be 2.34 kcal mol1. The value for the interaction energy of CO2 from a similar quality computational The electrostatic potential maps of the CO2 and the funcmethod and the position of CO2 above the ring is in very good tionalized benzenes are presented in Figure 2. Due to the charge separation in CO2, the electrostatic potential map is agreement with previous studies. The interaction energy of this system is the lowest that we found from our calculations separated in two electron rich regions (around the O atoms) with the exception of the fluorine-functionalized benzene ring, and one electron poor region (around the C atom). This is which worsens the interaction energy. In this case, fluorine acts a direct proof of the quadruple nature of CO2. The intense of as an electron-withdrawing group that reduces the interaction the electron poor and the electron rich regions of the maps of the quadrupole with the p system of the ring. Additionally, correlate well with the corresponding binding energies of the there is no contribution from the interaction of the oxygen complexes that are formed. Taking into account the geomeatom of CO2 with the fluorine atom. Another exception is in tries of the complexes formed, it can be concluded that the inthe case of CO2 interacting with C6H5CN. In the optimized geteraction of the electron-poor carbon atom with the p system ometry, CO2 lies in the same plane as the ring, near the CN of benzene is quite important for the stabilization of the comgroup, as can be seen in Figure 1. The oxygen of the CO2 plex. In the case of NH2-, SO3H-, and COOH-functionalized benzenes, the optimized geometries that were obtained have forms a weak hydrogen bond with the nearby hydrogen atom a direct correlation with the electrostatic potential maps of the of the phenyl ring and acts as a Lewis base, whereas the carbon atom of the CO2 interacts with the cyanide group. The binding energy (2.46 kcal mol1) is relatively low if we consider the separation between the hydrogen of the ring and the oxygen of CO2 (2.57 ), but it can be attributed to repulsion forces acting between the nitrogen atom of the CN group and the oxygen atom of CO2. The above-mentioned results for the optimized geometries and the corresponding binding energies can be explained qualitatively if we examine the electrostatic potential maps and the electron-density redistribution plots. In electrostatic potential maps, the charge distribution (electron-rich and electron-poor regions) of a molecule can be visualized. The varying intensities of the electrostatic potential can be visualized by using a color Figure 2. Electrostatic potential maps of CO2 (1), C6H5F (2), C6H6 (3), C6H5NO2 (4), C6H5CN (5), C6H5CF3 (6), C6H5OH spectrum, in which electron- (7), C6H5PH2 (8), C6H5SH (9), C6H5CH3 (10), C6H5OC2H5 (11), C6H5SO2NH2 (12), C6H5C2N3 (13), C6H5COOH (14), poor regions that correspond to C6H5SO3H (15), and C6H5NH2 (16). 2014 Wiley-VCH Verlag GmbH & Co. KGaA, Weinheim
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CHEMPHYSCHEM ARTICLES two interacting fragments. For these systems, electrostatic interactions play a dominant role in the final geometry of the complex. The electrostatic potential maps also explain why the complex is more stabilized in the cases of SO3H and COOH. For aniline, the strong electrostatic potential of the nitrogen atom of the amino group dominates the interaction with the electron-poor carbon atom of CO2. Still, there is a significant contribution from the electrostatic potential of the phenyl ring for aniline. The difference in the binding energy of NH2- and CH2N3-functionalized benzenes can be attributed to the reduced electrostatic potential of the phenyl ring of the latter with respect to NH2. Electron-density redistribution plots were also visualized as bright and dark regions that represent areas which gain or miss electron density, respectively. These plots are presented in Figure 3 for all complexes studied herein. This type of plot has particular importance in the study of CO2 complexes because it is an easy way to visualize the polarization of CO2 and the Lewis acid–Lewis base interactions that are present in these complexes. It is evident from Figure 3 that the binding energy of CO2 is analogous to the extent of electron redistribution on CO2 in the complex. The largest electron density redistribution is observed for the CO2–C6H5SO3H complex, whereas CO2– C6H5F presents the lowest redistribution. For complexes with the highest binding energies, it can be observed that there is electron density loss (dark regions) on the C atom of CO2, and a relative electron density gain bright regions) coming from the functional group of the functionalized benzene (complexes 10 to 15 in Figure 3). This is very well visualized in the case of
www.chemphyschem.org the interaction of CO2 with aniline, in which the C atom of the CO2 lose electron density, whereas there is a gain in electron density over the N atom of the amino group. Additionally, if an acidic hydrogen atom is present in the functional group of the functionalized benzene, there is an extra stabilization effect on the complex due to the redistribution of electron density between that hydrogen atom and the nearby oxygen atom of the CO2. It is observed that there is an electron density loss near H atoms and an electron density gain around the nearest O atom of the CO2. This effect can be clearly seen for the interaction of CO2 with SO3H-, COOH-, and SO2NH2-functionalized benzenes (corresponding to complexes 14, 13, and 11, respectively, in Figure 3). The strength of the interaction between CO2 and the functionalized benzenes is reflected in the change in the O-C-O angle. Figure 4 presents a correlation diagram between the interaction energy and the change in the O-C-O angle. It can be clearly seen that as the interaction of CO2 increases, an almost linear increase in the distortion of the O-C-O angle from 1808 is observed. The lowest distortion was observed for NO2-functionalized benzene and the highest was observed for COOHfunctionalized benzene. To make a better estimation of the binding energies of CO2 in the best-performing complexes, a (34) extrapolation scheme was applied and the results are presented in Table 2. It can be seen that the binding energy of CO2 increased in all cases by almost 1 kcal mol1. As noted in previous works, it is important to apply an extrapolation scheme to have an accurate estimation of the binding energy.
3. Conclusion
Figure 3. Electron-density redistribution plots of C6H5F (1), C6H6 (2), C6H5NO2 (3), C6H5CN (4), C6H5CF3 (5), C6H5OH (6), C6H5PH2 (7), C6H5SH (8), C6H5CH3 (9), C6H5OC2H5 (10), C6H5SO2NH2 (11), C6H5C2N3 (12), C6H5COOH (13), C6H5SO3H (14), and C6H5NH2 (15) with CO2.
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In summary, ab initio calculations were performed to study the interactions of carbon dioxide with various substituted benzenes. CO2 can act both as Lewis acid or Lewis base. Qualitatively, the difference in binding energies of substituent benzenes can be explained by the electrostatic potential maps and the electronic-density redistribution plots. The binding energies for CO2 range from 2.3 to 4.3 kcal mol1. In all complexes, when the electronic-density redistribution of the carbon atom of CO2 is more intense, a higher binding energy was found. In most of the studied cases, the CO2 molecule preferred a position on top of the ring and near the benzene functional groups. These geometrical configurations lead to a maximum of p···quadrupole interactions between the aroChemPhysChem 2014, 15, 905 – 911
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www.chemphyschem.org tionalized benzene. These stationary points were verified as minima by applying numerical frequency calculations.
Figure 4. Plot of the binding energy as a function of the O-C-O angle. A line was added to guide the eye.
To obtain a better estimation for the interaction energy for the structures that led to the highest interaction energies, an extrapolation scheme to approach the complete basis set (CBS) limit was applied. We performed single-point energy calculations on the structures that were optimized with RI-MP2/def2-TZVPP by enlarging the basis set. Single-point energy calculations were performed with Dunning’s augmented correlation consisted basis sets aug-ccpVXZ,[61] with X = 3 and X = 4. The MP2 correlation energies were extrapolated to the CBS limit by applying the two point Helgaker extrapolation scheme,[62] as can be seen in Equation (1): MP2 ECBS ðXY Þ ¼
Table 2. Interaction energies calculated by using a RI-MP2/def2-TZVPP, RI-MP2/aug-cc-pVTZ, RI-MP2/aug-ccpVQZ, and MP2/(34) extrapolation scheme. All energies were calculated by including counterpoise correction.
C6H5NH2–CO2 C6H5SO2NH2–CO2 C6H5CH2N3–CO2 C6H5COOH–CO2 C6H5SO3H–CO2
RI-MP2/def2-TZVPP [kcal mol1]
RI-MP2/aug-cc-pVTZ [kcal mol1]
RI-MP2/aug-cc-pVQZ [kcal mol1]
MP2(34) [kcal mol1]
3.34 3.48 3.62 3.65 4.27
3.89 4.26 4.37 4.24 4.98
4.04 4.37 4.52 4.44 5.23
4.14 4.46 4.63 4.55 5.36
matic ring and the CO2 molecule and an interaction between the functional group and carbon dioxide. The weak hydrogen bonding between the oxygen atom of CO2 and hydrogen atom of the substituent is important because it provides further stabilization of the complexes and enhances the binding energy as in cases of C6H5SO2NH2, C6H5COOH, C6H5SO3H, and C6H5NH2. In addition it is worth noting that substituents that are electron acceptors (F, NO2, CN, CF3) have weak binding energies for CO2 whereas substituent that are electron donors (OH, CH3, OC2H5, COOH, NH2) have larger CO2 binding energies, which enforces with the impact of hydrogen bonding.
Computational Section Second-order Møller–Plesset (MP2) perturbation theory in the resolution of identity (RI) approximation[57] was applied to our calculations along with the def2-TZVPP[58] basis set and the corresponding MP2 optimized auxiliary basis set for the RI approximation.[59] The SCF [self-consistent field Hartree–Fock (HF)] convergence criterion was set at 108 au. All calculations were performed by using the TURBOMOLE program.[60] The choice of post-HF Møller–Plesset perturbation method was based on its ability to take into account the contributions to the binding energy that arise from dispersion forces, with no extreme additional computational cost. The dispersion forces are dominant in nearly all cases where weak physisorption is involved. To locate the energetically most stable structure, several initial configurations were created by changing the position and orientation of the CO2 around each functionalized benzene. The initial configurations were generated by taking into account the complexity of the functional group and the corresponding electrostatic potential map for each functionalized benzene. Geometry optimization was performed for all initial configurations without any symmetry constraints and several stationary points were found (see Figure S1 in the Supporting Information) for each func 2014 Wiley-VCH Verlag GmbH & Co. KGaA, Weinheim
EXMP2 X 3 EYMP2 Y 3 X3 Y3
ð1Þ
in which X and Y are the cardinal numbers of the two basis sets used. The extrapolated MP2 correlation energies were added to the HF energy obtained from the largest basis set (aug-cc-pVQZ).
For all calculations, the binding energy of the complexes was calculated by using the supermolecule approach. The counterpoise correction for the basis-set superposition error (BSSE) was applied to correct the calculated binding energies according to the formula proposed by Boys and Bernardi.[63] The equation for the calculation of the binding energy is defined as shown in Equation (2): ghost ghost BE ¼ Ecomlpex Eorganic ECO þ DEdeform 2
ð2Þ
in which Ecomplex is the total energy of the complex, Eorganic, ghost is the energy of the organic molecule calculated at the complex geometry in the presence of the ghost basis of the CO2 molecule, E ghost CO2 is the energy of CO2 calculated at the complex geometry in the presence of the ghost basis of the organic molecule, and DEdeform is the deformation energy, defined as the difference between the isolated interacting molecules (CO2 and each one of the functionalized benzenes) in the complex geometry and in their optimized structures. The correction for the BSSE has been proven essential because this error is critical for nonbonding interactions. In an effort to understand and explain the interactions between CO2 and the functionalized benzenes, electrostatic potential maps were generated by mapping the electrostatic potentials onto the surfaces of molecular electron density by using gOpenMol.[64] The color of maps indicate regions of high and low electrostatic potential depending the density of electrons in these regions. The electrostatic potential in all structures ranged from + 0.03 to 0.03 hartree e1; blue signifies a value greater than or equal to the maximum in positive potential and red signifies a value greater than or equal to the maximum in negative potential. Electronic-density redistribution plots were also computed for all organic molecules interacting with CO2, according to the following equation: DD = D(organic···CO2)D(organic)D(CO2). The density of each monomer was evaluated at the MP2/def2-TZVPP level at the complex geometry and in the presence of the ghost basis functions of the other monomer. Densities were plotted with a contour value of 0.001 e Bohr3. Bright and dark regions represent areas that gain or lose electron density. Mathematical operations on the ChemPhysChem 2014, 15, 905 – 911
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CHEMPHYSCHEM ARTICLES densities and the plots were done by using module “contman” in gOpenMol.
Acknowledgements The present study was co-financed by the European Union (European Social Fund (ESF)) and Greek national funds through the Operational Program “Education and Lifelong Learning” of the National Strategic Reference Framework (NSRF) Research Funding Programs: Heracleitus II and Thales, investing in knowledge society through the European Social Fund. Keywords: ab initio calculations · CO2 capture · donor– acceptor systems · metal–organic frameworks · weak interactions [1] B. Metz, O. Davidson, H. de Coninck, M. Loos, L. Meyer, Carbon Dioxide Capture and Storage, Cambridge University Press, New York, 2005. [2] M. M. Abu-Khader, Energy Sources Part A 2006, 28, 1261. [3] M. Radosz, X. Hu, K. Krutkramelis, Y. Shen, Ind. Eng. Chem. Res. 2008, 47, 3783. [4] M. Smiglak, A. Metlen, R. D. Rogers, Acc. Chem. Res. 2007, 40, 1182. [5] S. Choi, J. H. Drese, C. W. Jones, ChemSusChem 2009, 2, 796. [6] P. J. E. Harlick, A. Sayari, Ind. Eng. Chem. Res. 2006, 45, 3248. [7] B. Feng, H. An, E. Tan, Energy Fuels 2007, 21, 426. [8] E. Ochoa-Fernndez, H. K. Rusten, H. A. Jakobsen, M. Ronning, A. Holmen, D. Chen, Catal. Today 2005, 106, 41. [9] Y. S. Cheng, Q. L. Huang, M. Eic, B. J. Balcom, Langmuir 2005, 21, 4376. [10] I. Deroche, L. Gaberova, G. Maurin, P. Llewellyn, M. Castro, P. Wright, Adsorption 2008, 14, 207. [11] N. Konduru, P. Lindner, N. M. Assaf-Anad, AIChE J. 2007, 53, 3137. [12] J. Pawlesa, A. Zukal, J. Cejka, Adsorption 2007, 13, 257. [13] R. T. Yang, Adsorbents: Fundamentals and Applications, Wiley, Hoboken, 2003. [14] M. C. B. Lpez, A. Martnez-Alonso, J. M. D. Tascn, Carbon 2000, 38, 1177. [15] F. Dreisbach, R. Staudt, J. U. Keller, Adsorption 1999, 5, 215. [16] P. K. Thallapally, P. B. McGrail, J. L. Atwood, C. Gaeta, C. Tedesco, P. Neri, Chem. Mater. 2007, 19, 3355. [17] P. K. Thallapally, P. B. McGrail, S. J. Dalgarno, H. T. Schaef, J. Tian, J. L. Atwood, Nat. Mater. 2008, 7, 146. [18] S. R. Batten, R. Robson, Angew. Chem. Int. Ed. 1998, 37, 1460; Angew. Chem. 1998, 110, 1558. [19] G. Frey, Dalton Trans. 2009, 4400. [20] G. Frey, Chem. Soc. Rev. 2008, 37, 191. [21] R. A. Fischer, C. Will, Angew. Chem. Int. Ed. 2008, 47, 8164; Angew. Chem. 2008, 120, 8285. [22] D. Tanaka, S. Kitagawa, Chem. Mater. 2008, 20, 922. [23] S. Kitagawa, R. Matsuda, Coord. Chem. Rev. 2007, 251, 2490. [24] C. J. Kepert, Chem. Commun. 2006, 695 – 700. [25] U. Mueller, M. Schubert, F. Teich, H. Puetter, K. Schierle-Arndt, J. Pastre, J. Mater. Chem. 2006, 16, 626. [26] A. G. Wong-Foy, A. J. Matzger, O. M. Yaghi, J. Am. Chem. Soc. 2006, 128, 3494. [27] A. R. Millward, O. M. Yaghi, J. Am. Chem. Soc. 2005, 127, 17998. [28] H. Furukawa, N. Ko, Y. B. Go, N. Aratani, S. B. Choi, E. Choi, A. O. Yazaydin, R. Q. Snurr, M. O’Keeffe, J. Kim, O. M. Yaghi, Science 2010, 329, 424 – 428. [29] J. W. Yoon, S. H. Jhung, Y. K. Hwang, S. M. Humphrey, P. T. Wood, J. S. Chang, Adv. Mater. 2007, 19, 1830.
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Ab Initio Study of the Adsorption of CO2 on Functionalized Benzenes Maria G. Frysali, Emmanuel Klontzas, and George. E. Froudakis*[a] The interaction of carbon dioxide with a series of functionalized aromatic molecules was studied by using quantum mechanical methods (MP2), to examine the effect of the substituent on the adsorption of CO2. Several different initial configurations of CO2 were taken into account for each functionalized benzene to locate the energetically most favorable configuration. To get a better estimation of the binding energies, we applied an extrapolation scheme to approach the complete basis set. CH2N3-, COOH-, and SO3H-functionalized benzenes were
found to have the strongest interaction with CO2, and the corresponding binding energies were calculated to be 3.62, 3.65, and 4.3 kcal mol1, respectively. Electrostatic potential maps of the functionalized benzenes and electron redistribution density plots of the complexes were also created to get a better insight into the nature of the interaction of CO2 with the functionalized benzenes. The functional groups that were examined can be potentially incorporated in organic bridging molecules that connect the inorganic corners in MOF.
1. Introduction Carbon dioxide is the fourth most abundant gas that is naturally present in the Earth’s atmosphere. Over the past two centuries its concentration in the atmosphere has greatly increased, mainly due to human activities, such as fossil fuel burning. Today, it is considered to be one of the greenhouse gases that contribute to the Earth’s global warming phenomenon. For the next few years, our energy demands will be satisfied by using carbon-based fuels as energy resources. To reduce the amount of CO2 that is released into the atmosphere, the global community is trying to find efficient ways to reduce the amount of CO2 released by using either direct capture from the air or from the resources from which it is produced, such as power plants or combustion engines. The whole process of limiting CO2 emissions due to human activities in the atmosphere is known as the carbon dioxide capture and storage (CCS) procedure.[1] This procedure includes a group of technologies for the capture of CO2 from power plants, followed by compression and transport to permanent storage sites. The capture process comprises the main cost of CCS schemes and represents approximately two thirds of the total cost. There are three processes that are used to capture CO2, known as post-combustion, pre-combustion, and oxyfuel processes. A diverse range of materials with different properties is used for capturing CO2 in these processes. Post-combustion methods have been considered to be less energy-de[a] M. G. Frysali, Dr. E. Klontzas, Prof. G. E. Froudakis Department of Chemistry University of Crete POBOX 2208 71003 Voutes Heraklion (Greece) Fax: (+ 30) 2810545001 E-mail: [email protected] Supporting Information for this article is available on the WWW under http://dx.doi.org/10.1002/cphc.201300952.
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manding methods and a lot of efforts have been devoted to finding suitable materials for this process. Many materials have been investigated for CO2 adsorption. Examples of these materials include physical absorbents,[2] ionic liquids,[3, 4] carbonbased sorbents,[5, 6] metal oxides,[7] hydrotalcite-like[8] compounds, zeolites,[2, 9–15] organic solids,[16, 17] and metal–organic frameworks. In the chemical industry, the separation of CO2 from mixtures with natural gas and flue gas is tremendously important. Natural gas is an alternative substitute for environmentally unfriendly fossil fuels. However, impurities in natural gas, such as CO2, reduce the calorie content and removal of CO2 from natural gas is expensive. For this reason, significant effort has been made in recent years to improve CO2 capture. Although a variety of technologies and methods have been developed, the discovery of new materials with high selectivity toward CO2 in a mixture of gases and high capacities remains a challenge. Of the different types of materials studied for this purpose, metal–organic frameworks (MOFs), which are organic–inorganic hybrid materials composed of metal ions or clusters interconnected through an organic linker, have attracted wide scientific interest owing to their enormous structural and chemical diversity and their potential applications in gas storage, ion exchange, molecular separation, and heterogeneous catalysis.[18–25] Recently, attention has turned to experimental and computational screening studies of different materials to assess CO2 adsorption. A large variety of MOFs, such as IRMOF1, IRMOF3, IRMOF6, IRMOF11, MOF2, MOF74, MOF177, MOF505, and [Cu3(BTC)2], were experimentally examined for their structural and porous attributes towards CO2 adsorption.[26] MOF-177, with a surface area (SA) of 4500 m2 g1, exhibits the highest capacity for CO2 and takes up 33.5 mol kg1 at 32 bar and
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298 K,[27] whereas MOF200 and MOF210[28] achieved CO2 caRegarding previous computational studies,[43–46] density func1 pacities of 54.5 mol kg at 50 bar and 298 K. tional theory (DFT) and post-Hartree–Fock methods[47–51] have At high pressures, CO2 capacities depend on the pore been applied to study many properties of these MOF materials. Traditional DFT functionals fail to evaluate the dispersion intervolume and surface areas[29, 30] of the MOFs. The increase in action, which plays an important role in weakly interacting pore volume and surface area enhance the storage capacity. In complexes, crystal packing of organic molecules, and three-dicontrast, CO2 capacities depend on the heats of adsorption for mensional structures of biological systems. Another disadvantCO2. Many studies have tried to improve the selectivity toward age of this method is the incorrect evaluation of polarization CO2 through polar functional groups, pore-size control,[31–33] and hyperpolarization abilities of large p-conjugated moleopen metal sites,[34–36] or with the introduction of alkali–metal cules.[52] In recent years, DFT and ab initio calculations have cations.[37] been used to study interaction energies and adsorption sites There are studies that indicate that MOFs with highly polar between MOFs and CO2.[53–56] Herein, we used ab initio methligands enhance the storage capacity for CO2. An example is bio-MOF-11, which presents high CO2 uptake at 298 K and very ods to study organic linkers that could improve the CO2 adminor N2 adsorption and is based on the Lewis basic sites of sorption capacity in MOFs. adenine, which includes an amino group and pyrimidine nitrogen atoms with strong affinity for CO2.[31] Another example is 2. Results and Discussion amine-functionalized MIL-53(Al). Amine increases the CO2/CH4 separation factor (from 5 to 60) compared with non-functionalOptimized structures of CO2 with the functionalized benzenes ized MIL-53.[38] Concerning the pore-size effect, there are severcan be seen in Figure 1 and the corresponding binding eneral MOFs with high selectivity for CO2 over N2 or CH4 due to kigies are reported in Table 1. Optimized geometries of the netic separation effect[30, 39] and other MOFs with indicated sestructures that correspond to the energetically most favorable lectivity due to a molecular sieving effect.[29, 40] Open metal structure were located and verified after performing geometry optimizations from different starting geometries. From sites have been also studied for CO2 capture and separation. Figure 1, it can be seen that the optimized structures can be For example, CO2/CH4 selectivities between carborane-based divided into two categories, that is, CO2 can be located above MOFs are compared with and without open metal sites and the benzene ring or it can be located over the functional the results suggested that high selectivity for CO2 over CH4 group. The binding energies calculated at the RI-MP2/def2(17) are obtained.[41] TZVPP level, reported in Table 1, range from 2.27 to In materials that capture carbon dioxide, such as MOFs, the adsorption of CO2 in the pores of the materials plays a major 4.27 kcal mol1 for CO2–C6H5F and CO2–C6H5SO3H, respectively. role. The knowledge of the nature of the interaction of CO2 The largest interaction energy was found when CO2 interactwith the building fragments that form the material is vital in understanding the adsorption process in existing materials and ed with C6H5SO3H, for which the interaction was calculated to in designing new materials for enhanced carbon dioxide capbe 4.27 kcal mol1. In this optimized structure, CO2 is not lo[42] ture. The nature of the interaction of CO2 with molecules has been attributed to the permanent quadrupole moment of CO2. CO2 can act as a Lewis acid or as a Lewis base. It acts as Lewis acid in the presence of electron-rich atoms on substituted benzenes, through interaction with the (electron-deficient) C atom of CO2, or as Lewis base in the presence of electrondeficient atoms of substituted benzenes through interaction with the (electron-rich) O atom of CO2. As has been noted in previous studies,[55, 56] dispersion forces and electrostatic interactions contribute almost equally to the binding energy of CO2 and, in some cases, dispersion forces have a more dominant Figure 1. Optimized structures of C6H5F (1), C6H6 (2), C6H5NO2 (3), C6H5CN (4), C6H5CF3 (5), C6H5OH (6), C6H5PH2 (7), role in the binding energy of the C6H5SH (8), C6H5CH3 (9), C6H5OC2H5 (10), C6H5SO2NH2 (11), C6H5C2N3 (12), C6H5COOH (13), C6H5SO3H (14), and molecular systems. C6H5NH2 (15) with CO2. 2014 Wiley-VCH Verlag GmbH & Co. KGaA, Weinheim
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Table 1. Binding energies (Ecp) and the corresponding O-C-O angles for all CO2 complexes calculated at the RI-MP2/def2-TZVPP level of theory. All energies were calculated by including a counterpoise correction.
C6H5F–CO2 C6H6–CO2 C6H5NO2–CO2 C6H5CN–CO2 C6H5CF3–CO2 C6H5OH–CO2 C6H5PH2–CO2 C6H5SH–CO2 C6H5CH3–CO2 C6H5OC2H5–CO2 C6H5NH2–CO2 C6H5SO2NH2–CO2 C6H5CH2N3–CO2 C6H5COOH–CO2 C6H5SO3H–CO2
Ecp [kcal mol1]
O-C-O angle [8]
2.27 2.34 2.42 2.46 2.55 2.72 2.80 2.92 2.92 3.22 3.34 3.48 3.62 3.65 4.27
179.161 179.024 179.472 178.844 179.418 179.033 179.140 179.132 178.851 178.791 178.484 178.600 178.947 177.988 178.360
cated above the benzene ring, but close to the SO3H functional group. The complex is stabilized by the simultaneous interaction of the carbon atom of CO2 with the oxygen atom of the sulfonate group and the interaction the oxygen of CO2 with the hydrogen atom of the sulfonate group. The CCO2OSO3H and OCO2HSO3H contacts were measured to be 2.90 and 1.98 , respectively. As mentioned above in the Introduction, both the carbon and oxygen atoms of the CO2 can contribute to the stabilization of the complex by acting as a Lewis acid and as a Lewis base, respectively. In this case, the electron-deficient carbon atom of CO2 acts as Lewis acid when interacting with the oxygen atom of the sulfonate group. The oxygen atom of CO2 acts as a Lewis base and forms a hydrogen bond with the electron-deficient hydrogen atom of the sulfonate group, as indicated by the short OCO2HSO3H contact. The next highest binding energy was calculated to be 3.65 kcal mol1 for the interaction of CO2 with C6H5COOH. As in the previous compound, the oxygen of the carboxylic group act as a Lewis base to the electron-deficient carbon atom of the CO2, whereas the hydrogen atom acts as a Lewis acid. The CCO2OCOOH and OCO2HCOOH contacts were measured to be 2.85 and 2.03 , respectively, which are similar to the corresponding values in C6H5SO3H. CO2 is oriented in the same plane, which can be formed from the atoms of the carboxylic group. CO2 interacts with C6H5CH2N3 with a binding energy of 3.62 kcal mol1, which is similar to the binding energy found for the carboxylic group. In this case, CO2 is oriented over CH2N3 in such a way that it can interact with both the nitrogen atom of the azide group and the hydrogen atom of the CH2. Based on the optimized geometry that was obtained from our calculations and the corresponding separations between the atoms of CO2 and the atoms of the CH2N3 group, N3 act as a Lewis base and the formed complex is further stabilized by the formation of a weak hydrogen bond between the oxygen atom of CO2 and the neighboring hydrogen atom of CH2. This is supported by the corresponding separations between the nitrogen atom of N3 with the carbon atom of CO2 (3.09 ) and that of the oxygen atom of the CO2 and the hydrogen atom of CH2 2014 Wiley-VCH Verlag GmbH & Co. KGaA, Weinheim
(2.89 ). The value of 3.09 for the CN contact is a bit longer than the corresponding separation in molecular systems that contain nitrogen atoms that interact with CO2 ( 2.90 ). The binding energy that was obtained for C6H5CH2N3 is larger than the value (2.56 kcal mol1) that was calculated by Vitillo et al.[55] for C6H5N3 with a similar quality of computational methods. This can be attributed to the existence of the weak hydrogen bond that we mentioned previously. The interaction energy of CO2 with C6H5SO2NH2 was calculated to be 3.48 kcal mol1. Again, we have an electron donor– acceptor mechanism for the stabilization of this complex. One of the oxygen atoms of CO2 forms a hydrogen bond with a hydrogen of the NH2 group and the corresponding separation was calculated at 2.18 . The carbon atom of the CO2 acts as a Lewis acid and interacts with the oxygen atom of the SO2NH2 group. The contact between the carbon atom and the oxygen atom is 2.95 , similar to the value found for the CO2– C6H5SO3H complex. The next lowest binding energy was found for the interaction of CO2 with C6H5NH2. The interaction energy for this system was calculated to be 3.34 kcal mol1, in agreement with the value of 3.27 kcal mol1 that was previously found by Vitillo et al.[55] CO2 prefers a location above the N atom of the amino group due to the electron donor–acceptor mechanism between the C atom of CO2 and the N atom of the amino group. We also investigated the interaction of CO2 with C6H5OC2H5. The interaction energy was calculated to be 3.22 kcal mol1. The geometrical configuration of the CO2 is characterized by the formation of two hydrogen bonds between OCO2 and HOC2H5, with a contact of 2.67 for OCO2H and a contact of 2.93 between OCO2 and the H atom of benzene. The C atom of CO2 interacts with the O atom of OC2H5, as can be deduced from their interatomic separation (2.91 ). In the optimized geometries found for CO2 with the other functionalized benzenes, we observed that the CO2 is oriented in a parallel mode above the plane of the benzene ring, with an exception in the case of C6H5CN. In all cases except the latter, CO2 is displaced from the center of mass of the benzene ring, which indicates that there is an additional interaction between the oxygen atoms of the CO2 with the atoms of the polar groups on top of the known interaction between the quadrupole moment of CO2 and the p system of benzene. This phenomenon has also been mentioned in previous studies on the interaction of CO2 with molecular systems that contain aromatic rings. For the interaction of CO2 with C6H5CH3, the calculated value for the interaction was found to be 2.92 kcal mol1. In this case, the methyl group acts as an electron donor to the benzene ring and enhances the interaction of the quadrupole with the p system of the ring. Furthermore, there is an additional contribution due to the interaction between the oxygen of CO2 and the hydrogen from the methyl group, which is reflected in the short contact of 2.66 between these two atoms. The same interaction energy was found also for the interaction of CO2 with C6H5SH. This functional group acts as an electron acceptor and weakens the p system of the ring. The geometry of the optimized structure is stabilized through the interaction of the hydrogen of the SH group with the polarized oxygen of ChemPhysChem 2014, 15, 905 – 911
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a high potential value are represented in red and electron-rich CO2 with an OCO2HSH contact of 2.96 . CO2 interacts with regions that correspond to a low potential are represented C6H5PH2 and C6H5OH with interaction energies of 2.80 and with blue. Electrostatic potential maps can be useful in explain2.67 kcal mol1, respectively. In the case of C6H5OH, CO2 is loing and predicting the geometry of the molecular system and cated on the top of the benzene ring, which means that the the nature of the interactions between two molecules, espemost stable geometry is when CO2 interacts not only with the cially when electrostatic interactions are important in describformation of a weak hydrogen bond between OCO2 and HOH ing their interaction. The stability of the systems that were exbut with the simultaneous interaction of CO2 with the p amined herein comes from the existence of electrostatic intersystem of the ring. The separation between the oxygen atom actions due to the charged nature of CO2 and the functionalof CO2 and the hydrogen atom of the hydroxyl group was measured to be 2.98 . ized benzenes on top of the contribution due to the existence For the other functionalized benzenes that were studied, the of dispersion forces. In fact, previous studies have shown that interaction energies with CO2 ranged from 2.27 to 2.55 kcal dispersion forces and electrostatic interactions contribute almost equally to the binding energy of CO2 and, in some mol1. The interaction energy of CO2 with benzene, for which the location of CO2 is distorted from the center of mass of the cases, dispersion forces have more a dominant role in the binding energy of the molecular systems. ring, was calculated to be 2.34 kcal mol1. The value for the interaction energy of CO2 from a similar quality computational The electrostatic potential maps of the CO2 and the funcmethod and the position of CO2 above the ring is in very good tionalized benzenes are presented in Figure 2. Due to the charge separation in CO2, the electrostatic potential map is agreement with previous studies. The interaction energy of this system is the lowest that we found from our calculations separated in two electron rich regions (around the O atoms) with the exception of the fluorine-functionalized benzene ring, and one electron poor region (around the C atom). This is which worsens the interaction energy. In this case, fluorine acts a direct proof of the quadruple nature of CO2. The intense of as an electron-withdrawing group that reduces the interaction the electron poor and the electron rich regions of the maps of the quadrupole with the p system of the ring. Additionally, correlate well with the corresponding binding energies of the there is no contribution from the interaction of the oxygen complexes that are formed. Taking into account the geomeatom of CO2 with the fluorine atom. Another exception is in tries of the complexes formed, it can be concluded that the inthe case of CO2 interacting with C6H5CN. In the optimized geteraction of the electron-poor carbon atom with the p system ometry, CO2 lies in the same plane as the ring, near the CN of benzene is quite important for the stabilization of the comgroup, as can be seen in Figure 1. The oxygen of the CO2 plex. In the case of NH2-, SO3H-, and COOH-functionalized benzenes, the optimized geometries that were obtained have forms a weak hydrogen bond with the nearby hydrogen atom a direct correlation with the electrostatic potential maps of the of the phenyl ring and acts as a Lewis base, whereas the carbon atom of the CO2 interacts with the cyanide group. The binding energy (2.46 kcal mol1) is relatively low if we consider the separation between the hydrogen of the ring and the oxygen of CO2 (2.57 ), but it can be attributed to repulsion forces acting between the nitrogen atom of the CN group and the oxygen atom of CO2. The above-mentioned results for the optimized geometries and the corresponding binding energies can be explained qualitatively if we examine the electrostatic potential maps and the electron-density redistribution plots. In electrostatic potential maps, the charge distribution (electron-rich and electron-poor regions) of a molecule can be visualized. The varying intensities of the electrostatic potential can be visualized by using a color Figure 2. Electrostatic potential maps of CO2 (1), C6H5F (2), C6H6 (3), C6H5NO2 (4), C6H5CN (5), C6H5CF3 (6), C6H5OH spectrum, in which electron- (7), C6H5PH2 (8), C6H5SH (9), C6H5CH3 (10), C6H5OC2H5 (11), C6H5SO2NH2 (12), C6H5C2N3 (13), C6H5COOH (14), poor regions that correspond to C6H5SO3H (15), and C6H5NH2 (16). 2014 Wiley-VCH Verlag GmbH & Co. KGaA, Weinheim
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CHEMPHYSCHEM ARTICLES two interacting fragments. For these systems, electrostatic interactions play a dominant role in the final geometry of the complex. The electrostatic potential maps also explain why the complex is more stabilized in the cases of SO3H and COOH. For aniline, the strong electrostatic potential of the nitrogen atom of the amino group dominates the interaction with the electron-poor carbon atom of CO2. Still, there is a significant contribution from the electrostatic potential of the phenyl ring for aniline. The difference in the binding energy of NH2- and CH2N3-functionalized benzenes can be attributed to the reduced electrostatic potential of the phenyl ring of the latter with respect to NH2. Electron-density redistribution plots were also visualized as bright and dark regions that represent areas which gain or miss electron density, respectively. These plots are presented in Figure 3 for all complexes studied herein. This type of plot has particular importance in the study of CO2 complexes because it is an easy way to visualize the polarization of CO2 and the Lewis acid–Lewis base interactions that are present in these complexes. It is evident from Figure 3 that the binding energy of CO2 is analogous to the extent of electron redistribution on CO2 in the complex. The largest electron density redistribution is observed for the CO2–C6H5SO3H complex, whereas CO2– C6H5F presents the lowest redistribution. For complexes with the highest binding energies, it can be observed that there is electron density loss (dark regions) on the C atom of CO2, and a relative electron density gain bright regions) coming from the functional group of the functionalized benzene (complexes 10 to 15 in Figure 3). This is very well visualized in the case of
www.chemphyschem.org the interaction of CO2 with aniline, in which the C atom of the CO2 lose electron density, whereas there is a gain in electron density over the N atom of the amino group. Additionally, if an acidic hydrogen atom is present in the functional group of the functionalized benzene, there is an extra stabilization effect on the complex due to the redistribution of electron density between that hydrogen atom and the nearby oxygen atom of the CO2. It is observed that there is an electron density loss near H atoms and an electron density gain around the nearest O atom of the CO2. This effect can be clearly seen for the interaction of CO2 with SO3H-, COOH-, and SO2NH2-functionalized benzenes (corresponding to complexes 14, 13, and 11, respectively, in Figure 3). The strength of the interaction between CO2 and the functionalized benzenes is reflected in the change in the O-C-O angle. Figure 4 presents a correlation diagram between the interaction energy and the change in the O-C-O angle. It can be clearly seen that as the interaction of CO2 increases, an almost linear increase in the distortion of the O-C-O angle from 1808 is observed. The lowest distortion was observed for NO2-functionalized benzene and the highest was observed for COOHfunctionalized benzene. To make a better estimation of the binding energies of CO2 in the best-performing complexes, a (34) extrapolation scheme was applied and the results are presented in Table 2. It can be seen that the binding energy of CO2 increased in all cases by almost 1 kcal mol1. As noted in previous works, it is important to apply an extrapolation scheme to have an accurate estimation of the binding energy.
3. Conclusion
Figure 3. Electron-density redistribution plots of C6H5F (1), C6H6 (2), C6H5NO2 (3), C6H5CN (4), C6H5CF3 (5), C6H5OH (6), C6H5PH2 (7), C6H5SH (8), C6H5CH3 (9), C6H5OC2H5 (10), C6H5SO2NH2 (11), C6H5C2N3 (12), C6H5COOH (13), C6H5SO3H (14), and C6H5NH2 (15) with CO2.
2014 Wiley-VCH Verlag GmbH & Co. KGaA, Weinheim
In summary, ab initio calculations were performed to study the interactions of carbon dioxide with various substituted benzenes. CO2 can act both as Lewis acid or Lewis base. Qualitatively, the difference in binding energies of substituent benzenes can be explained by the electrostatic potential maps and the electronic-density redistribution plots. The binding energies for CO2 range from 2.3 to 4.3 kcal mol1. In all complexes, when the electronic-density redistribution of the carbon atom of CO2 is more intense, a higher binding energy was found. In most of the studied cases, the CO2 molecule preferred a position on top of the ring and near the benzene functional groups. These geometrical configurations lead to a maximum of p···quadrupole interactions between the aroChemPhysChem 2014, 15, 905 – 911
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www.chemphyschem.org tionalized benzene. These stationary points were verified as minima by applying numerical frequency calculations.
Figure 4. Plot of the binding energy as a function of the O-C-O angle. A line was added to guide the eye.
To obtain a better estimation for the interaction energy for the structures that led to the highest interaction energies, an extrapolation scheme to approach the complete basis set (CBS) limit was applied. We performed single-point energy calculations on the structures that were optimized with RI-MP2/def2-TZVPP by enlarging the basis set. Single-point energy calculations were performed with Dunning’s augmented correlation consisted basis sets aug-ccpVXZ,[61] with X = 3 and X = 4. The MP2 correlation energies were extrapolated to the CBS limit by applying the two point Helgaker extrapolation scheme,[62] as can be seen in Equation (1): MP2 ECBS ðXY Þ ¼
Table 2. Interaction energies calculated by using a RI-MP2/def2-TZVPP, RI-MP2/aug-cc-pVTZ, RI-MP2/aug-ccpVQZ, and MP2/(34) extrapolation scheme. All energies were calculated by including counterpoise correction.
C6H5NH2–CO2 C6H5SO2NH2–CO2 C6H5CH2N3–CO2 C6H5COOH–CO2 C6H5SO3H–CO2
RI-MP2/def2-TZVPP [kcal mol1]
RI-MP2/aug-cc-pVTZ [kcal mol1]
RI-MP2/aug-cc-pVQZ [kcal mol1]
MP2(34) [kcal mol1]
3.34 3.48 3.62 3.65 4.27
3.89 4.26 4.37 4.24 4.98
4.04 4.37 4.52 4.44 5.23
4.14 4.46 4.63 4.55 5.36
matic ring and the CO2 molecule and an interaction between the functional group and carbon dioxide. The weak hydrogen bonding between the oxygen atom of CO2 and hydrogen atom of the substituent is important because it provides further stabilization of the complexes and enhances the binding energy as in cases of C6H5SO2NH2, C6H5COOH, C6H5SO3H, and C6H5NH2. In addition it is worth noting that substituents that are electron acceptors (F, NO2, CN, CF3) have weak binding energies for CO2 whereas substituent that are electron donors (OH, CH3, OC2H5, COOH, NH2) have larger CO2 binding energies, which enforces with the impact of hydrogen bonding.
Computational Section Second-order Møller–Plesset (MP2) perturbation theory in the resolution of identity (RI) approximation[57] was applied to our calculations along with the def2-TZVPP[58] basis set and the corresponding MP2 optimized auxiliary basis set for the RI approximation.[59] The SCF [self-consistent field Hartree–Fock (HF)] convergence criterion was set at 108 au. All calculations were performed by using the TURBOMOLE program.[60] The choice of post-HF Møller–Plesset perturbation method was based on its ability to take into account the contributions to the binding energy that arise from dispersion forces, with no extreme additional computational cost. The dispersion forces are dominant in nearly all cases where weak physisorption is involved. To locate the energetically most stable structure, several initial configurations were created by changing the position and orientation of the CO2 around each functionalized benzene. The initial configurations were generated by taking into account the complexity of the functional group and the corresponding electrostatic potential map for each functionalized benzene. Geometry optimization was performed for all initial configurations without any symmetry constraints and several stationary points were found (see Figure S1 in the Supporting Information) for each func 2014 Wiley-VCH Verlag GmbH & Co. KGaA, Weinheim
EXMP2 X 3 EYMP2 Y 3 X3 Y3
ð1Þ
in which X and Y are the cardinal numbers of the two basis sets used. The extrapolated MP2 correlation energies were added to the HF energy obtained from the largest basis set (aug-cc-pVQZ).
For all calculations, the binding energy of the complexes was calculated by using the supermolecule approach. The counterpoise correction for the basis-set superposition error (BSSE) was applied to correct the calculated binding energies according to the formula proposed by Boys and Bernardi.[63] The equation for the calculation of the binding energy is defined as shown in Equation (2): ghost ghost BE ¼ Ecomlpex Eorganic ECO þ DEdeform 2
ð2Þ
in which Ecomplex is the total energy of the complex, Eorganic, ghost is the energy of the organic molecule calculated at the complex geometry in the presence of the ghost basis of the CO2 molecule, E ghost CO2 is the energy of CO2 calculated at the complex geometry in the presence of the ghost basis of the organic molecule, and DEdeform is the deformation energy, defined as the difference between the isolated interacting molecules (CO2 and each one of the functionalized benzenes) in the complex geometry and in their optimized structures. The correction for the BSSE has been proven essential because this error is critical for nonbonding interactions. In an effort to understand and explain the interactions between CO2 and the functionalized benzenes, electrostatic potential maps were generated by mapping the electrostatic potentials onto the surfaces of molecular electron density by using gOpenMol.[64] The color of maps indicate regions of high and low electrostatic potential depending the density of electrons in these regions. The electrostatic potential in all structures ranged from + 0.03 to 0.03 hartree e1; blue signifies a value greater than or equal to the maximum in positive potential and red signifies a value greater than or equal to the maximum in negative potential. Electronic-density redistribution plots were also computed for all organic molecules interacting with CO2, according to the following equation: DD = D(organic···CO2)D(organic)D(CO2). The density of each monomer was evaluated at the MP2/def2-TZVPP level at the complex geometry and in the presence of the ghost basis functions of the other monomer. Densities were plotted with a contour value of 0.001 e Bohr3. Bright and dark regions represent areas that gain or lose electron density. Mathematical operations on the ChemPhysChem 2014, 15, 905 – 911
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CHEMPHYSCHEM ARTICLES densities and the plots were done by using module “contman” in gOpenMol.
Acknowledgements The present study was co-financed by the European Union (European Social Fund (ESF)) and Greek national funds through the Operational Program “Education and Lifelong Learning” of the National Strategic Reference Framework (NSRF) Research Funding Programs: Heracleitus II and Thales, investing in knowledge society through the European Social Fund. Keywords: ab initio calculations · CO2 capture · donor– acceptor systems · metal–organic frameworks · weak interactions [1] B. Metz, O. Davidson, H. de Coninck, M. Loos, L. Meyer, Carbon Dioxide Capture and Storage, Cambridge University Press, New York, 2005. [2] M. M. Abu-Khader, Energy Sources Part A 2006, 28, 1261. [3] M. Radosz, X. Hu, K. Krutkramelis, Y. Shen, Ind. Eng. Chem. Res. 2008, 47, 3783. [4] M. Smiglak, A. Metlen, R. D. Rogers, Acc. Chem. Res. 2007, 40, 1182. [5] S. Choi, J. H. Drese, C. W. Jones, ChemSusChem 2009, 2, 796. [6] P. J. E. Harlick, A. Sayari, Ind. Eng. Chem. Res. 2006, 45, 3248. [7] B. Feng, H. An, E. Tan, Energy Fuels 2007, 21, 426. [8] E. Ochoa-Fernndez, H. K. Rusten, H. A. Jakobsen, M. Ronning, A. Holmen, D. Chen, Catal. Today 2005, 106, 41. [9] Y. S. Cheng, Q. L. Huang, M. Eic, B. J. Balcom, Langmuir 2005, 21, 4376. [10] I. Deroche, L. Gaberova, G. Maurin, P. Llewellyn, M. Castro, P. Wright, Adsorption 2008, 14, 207. [11] N. Konduru, P. Lindner, N. M. Assaf-Anad, AIChE J. 2007, 53, 3137. [12] J. Pawlesa, A. Zukal, J. Cejka, Adsorption 2007, 13, 257. [13] R. T. Yang, Adsorbents: Fundamentals and Applications, Wiley, Hoboken, 2003. [14] M. C. B. Lpez, A. Martnez-Alonso, J. M. D. Tascn, Carbon 2000, 38, 1177. [15] F. Dreisbach, R. Staudt, J. U. Keller, Adsorption 1999, 5, 215. [16] P. K. Thallapally, P. B. McGrail, J. L. Atwood, C. Gaeta, C. Tedesco, P. Neri, Chem. Mater. 2007, 19, 3355. [17] P. K. Thallapally, P. B. McGrail, S. J. Dalgarno, H. T. Schaef, J. Tian, J. L. Atwood, Nat. Mater. 2008, 7, 146. [18] S. R. Batten, R. Robson, Angew. Chem. Int. Ed. 1998, 37, 1460; Angew. Chem. 1998, 110, 1558. [19] G. Frey, Dalton Trans. 2009, 4400. [20] G. Frey, Chem. Soc. Rev. 2008, 37, 191. [21] R. A. Fischer, C. Will, Angew. Chem. Int. Ed. 2008, 47, 8164; Angew. Chem. 2008, 120, 8285. [22] D. Tanaka, S. Kitagawa, Chem. Mater. 2008, 20, 922. [23] S. Kitagawa, R. Matsuda, Coord. Chem. Rev. 2007, 251, 2490. [24] C. J. Kepert, Chem. Commun. 2006, 695 – 700. [25] U. Mueller, M. Schubert, F. Teich, H. Puetter, K. Schierle-Arndt, J. Pastre, J. Mater. Chem. 2006, 16, 626. [26] A. G. Wong-Foy, A. J. Matzger, O. M. Yaghi, J. Am. Chem. Soc. 2006, 128, 3494. [27] A. R. Millward, O. M. Yaghi, J. Am. Chem. Soc. 2005, 127, 17998. [28] H. Furukawa, N. Ko, Y. B. Go, N. Aratani, S. B. Choi, E. Choi, A. O. Yazaydin, R. Q. Snurr, M. O’Keeffe, J. Kim, O. M. Yaghi, Science 2010, 329, 424 – 428. [29] J. W. Yoon, S. H. Jhung, Y. K. Hwang, S. M. Humphrey, P. T. Wood, J. S. Chang, Adv. Mater. 2007, 19, 1830.
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