Molecule Bonding in Lewis ... - ACS Publications

Article pubs.acs.org/JPCA

Directionality of Cation/Molecule Bonding in Lewis Bases Containing the Carbonyl Group Younes Valadbeigi*,† and Jean-François Gal‡ †

Department of Chemistry, Science and Research Branch, Islamic Azad University, Tehran, Iran Institut de Chimie de Nice, UMR 7272, Université Côte d’Azur, CNRS, 06108 Nice, France

‡

S Supporting Information *

ABSTRACT: Relationship between the CO−X+ (X = H, Li, Na, K, Al, Cu) angle and covalent characteristic of the X+−M (M = CH2O, CH3CHO, acetone, imidazol-2one (C2H2N2O), cytosine, γ-butyrolactone) was investigated, theoretically. The calculated electron densities ρ at the bond critical points revealed that the covalency of the M−X+ interaction depended on the nature of the cation and varied as H+ > Cu+ > Al+ > Li+ > Na+ > K+. The alkali cations tended to participate in electrostatic interactions and aligned with the direction of the molecule dipole or local dipole of CO group to form linear CO−X geometries. Because of overlapping with lonepair electrons of the sp2 carbonyl oxygen, the H+ and Cu+ formed a bent CO−X angle. Al+ displayed an intermediate behavior; the CO−Al angle was 180° in [CH2O/Al]+ (mainly electrostatic), but when the angle was bent (146°) under the effect of local dipole of an adjacent imine group in cytosine, the covalency of the CO−Al+ interaction increased. The CO−X angles in M/X+ adduct ions were scanned in different O−X bond lengths. It was found that the most favorable CO− X angle depended on the O−X bond length. This dependency was attributed to variation of covalent and electrostatic contributions with O−X distance. In addition, the structures of [CH2S/X]+ and [CH2Se/X]+ were studied, and only bent CS− X and CSe−X angles were obtained for all cations, although the dipole vectors of CH2S and CH2Se coincide with the CS and CSe bonds. The bending of the CS−X and CSe−X angles was attributed to the covalent characteristic of S−X and Se−X interactions due to high polarizability of S and Se atoms. covalent character of Al+ interactions.17 Similarly, in cytosine, Na+, Ag+, and Cu+ form chelate structures with O and N atoms of carbonyl and imine groups, while Al+ is attached only to the O atom of carbonyl group,18,19 which is not in accordance with the sequence of increasing covalency: Na+, Al+, Ag+, Cu+. Therefore, chelate structures are not always associated with electrostatic interactions. The Lewis bases containing N, P, O, and S atoms interact mainly with the cations via their heteroatoms. This interaction may be covalent via the electron lone pair(s) on the heteroatoms or electrostatic via cation/dipole interaction. On the one hand, it has been reported that the covalent character of the interactions of H+, CH3+, Cu+, and to some extent Al+, with Lewis bases containing these heteroatoms is considerable.13,20 On the other hand, the interactions of Li+, Na+, and K+ with the heteroatoms are mainly electrostatic ion/dipole interactions.20 For covalent interactions, the direction of the cation/molecule bond is determined by the axis of putative sp, sp2, and sp3 orbitals of lone pair electrons of heteroatoms.5 For a pure electrostatic interaction, the cation is aligned with the direction of the dipole moment of the molecule. For example,

1. INTRODUCTION Ion/molecule interactions are of primary importance not only in biology but also in environmental, atmospheric, and analytical chemistry and reaction mechanisms and have been well-studied in detail for different systems.1−5 Fundamental studies focused essentially on two aspects: energetic and structural properties of the systems involving ion/molecule interactions. The energetic studies of cation/molecule interactions in the gas phase using theoretical and experimental methods have resulted in large tabulated data of cation affinities and gas-phase cation basicities. 6−13 Some experimental techniques such as infrared (IR) spectroscopy can be used to investigate the structures of ion/molecule complexes; however, the interpretation of energetic and spectroscopic data is enhanced by computational chemistry methods.5,14,15 For complex molecules with several coordination centers, they may simultaneously interact with the cation (X+) via two or more sites to form chelate structures. The formation of chelate structures depends on the cation type or nature of molecule/X+ interaction. Generally, the ionization energy (IE) of X determines the nature of molecule/X+ interaction, so that the cations with larger IE tend to participate in more covalent interactions. Alkali metal cations are chelated in azines and azoles by two adjacent nitrogen atoms,16 while Al+ interacts with only one nitrogen atom, a fact that was attributed to © XXXX American Chemical Society

Received: May 10, 2017 Revised: August 23, 2017

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DOI: 10.1021/acs.jpca.7b04474 J. Phys. Chem. A XXXX, XXX, XXX−XXX

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The Journal of Physical Chemistry A

Figure 1. Optimized structures of CH2O and CH3CHO adduct ions of Li+, Na+, K+, Al+, Cu+, and H+. The vectors show the direction of dipole moments of neutral CH2O and CH3CHO.

compounds and 10 carboxylic acids.11,25 The results showed that the energies obtained by the B3LYP are more accurate than those obtained by the expensive methods such as CBS-Q, G4MP2, and W1BD.11,25 The nature of the cation/molecule interactions was investigated by quantum theory of atoms in molecules (QTAIM) using the AIM2000 software.26 The electron density ρ, its Laplacian ∇2ρ, and potential and kinetic electron energy densities (V(r) and G(r), respectively) at cation/molecule critical points CP, were used to compare the covalency of the interactions. The −G(r)/V(r) ratio for noncovalent interactions is greater than 1.27 Also, a larger value of ρ indicates more covalent nature.28 The sign of ∇2ρ may indicate the nature of interaction, so that for covalent interactions, the Laplacian is negative; however, this condition is not always satisfactory criterion.29 In addition, some authors claim that positive Laplacian and negative H(r) indicate partial covalency of the interaction, where H(r) = V(r) + G(r).29

in the adduct ions of H+ and Li+ with formaldehyde, the angles of CO−H and CO−Li are ∼120° and 180°, respectively.5 For amines, the axis of the lone-pair electrons on N approximately coincides with the direction of the molecule dipole moment;21 hence, the direction of the cations toward the molecule does not differ so much for electrostatic and covalent interactions. Conversely, molecules containing a carbonyl (CO) group, in which the dipole moment and sp2 lone pairs are aligned along different axes, are good candidates to study the effect of nature of cation/molecule interactions on the structures of adduct ions. In this work, the interactions of H+, Li+, Na+, K+, Al+, and Cu+ cations (X+) with carbonyl groups of simple (CH2O, CH3CHO, and (CH3)2CO) and more complex (imidazol-2one (C3H2N2O), cytosine, γ-butyrolactone) molecules are investigated. The nature of cation/molecule interactions on the CO−X angle in different O−X bond lengths is investigated. The study of the carbonyl group is extended to thioformaldehyde and selenoformaldehyde.

3. RESULTS AND DISCUSSION 3.1. Formaldehyde, Acetaldehyde, and Acetone. Figure 1 shows the optimized structures of cation adducts of formaldehyde (CH2O) and acetaldehyde (CH3CHO) as well as the O−X bond lengths and CO−X angles. The direction of the dipole moment of CH2O coincides with the direction of its CO bond, while a small deviation is observed for direction of the CH3CHO dipole moment. In CH2O and CH3CHO adduct ions, the angles CO−X+ for alkali cations are 180° and 175°, respectively, indicating mainly electrostatic interactions in the direction of the dipole moments. In the protonated molecules, the CO−H angle is ∼115° due to covalent interaction between H+ and lone-pair electrons of oxygen (sp2). In the Cu+ adduct ions with CH2O and CH3CHO, the CO−Cu angles are 143.7° and 139.5°, respectively, indicating that the covalent contribution is significant; however, covalency of O−Cu interaction is not as much as that of O−H. The cation affinities of CH2O and CH3CHO toward H+, Li+, Na+, K+, Al+, and Cu+ somewhat show the nature of the interactions (Table S1); that is, more covalent interactions have larger cation affinities. The nature of the O−X interactions fairly depends on the ionization

2. COMPUTATIONAL DETAILS The structures of the molecules and their cation/molecule adduct ions were optimized by B3LYP method using 6-311+ +G(d,p) basis set. Frequency calculations were performed at the same level of theory to obtain thermodynamic properties including proton affinities (PA), gas-phase basicities (GB), and cation affinities (CA) of the molecules at 298.15 K. The C O−X angles were scanned by the B3LYP/6-311++G(d,p) method for different O−X bond lengths. The other geometry parameters of the system were fixed during the angle scanning. All calculations were performed using the Gaussian 09 software.22 The accuracy of the B3LYP method has been assessed by several authors for different systems.11,23−25 Bryantsev et al.24 compared the accuracy of different functionals including B3LYP, X3LYP, M06-2X, and M06-L and reported that the B3LYP functional gives more accurate energies for systems including ion−molecules interaction. Also, we employed B3LYP, CBS-Q, G4MP2, and W1BD methods for calculation of proton and cation affinities of more than 60 organic B


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Figure 2. Potential energy curves obtained by scanning the CO−X angle of (a) [CH2O/Li]+, (b) [CH2O/K]+, (c) [CH2O/Al]+, and (d) [CH2O/ Cu]+ at different O−X bond lengths.

cation is an indication of the orbital overlap and an indirect measure of covalence. In the case of [CH2O/H]+ and [CH2O/ Cu]+ complexes, the covalent contribution overrides, and only bent CO−H and CO−Cu angles are obtained with different computational methods. Also, in the [CH2O/Li]+, [CH2O/Na]+, and [CH2O/K]+ complexes the electrostatic part is determinant factor, so that different methods predict only linear CO−X (X = Li, Na, K) geometries. In the [CH2O/ Al]+ complex, the electrostatic and covalent contributions are

energy (IE) of X. The calculated and experimental IE values for H, Li, Na, K, Al, and Cu were summarized in Table S2. The IE values decrease as H > Cu > Al > Li > Na > K. The observed trend is in accordance with the nature of the cation interactions; specifically, H+ and Cu+ mainly form covalent interactions, while alkali metal cations participate in electrostatic interactions. Furthermore, Mulliken charge distribution analysis (Figure S1) validates the covalent nature of H+ and Cu+ interactions. The amount of charge transferred to the C


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Figure 3. Potential energy curves obtained by scanning the CO−X angle of (a) [CH3CHO/Li]+, (b) [CH3CHO/K]+, (c) [CH3CHO/Al]+, and (d) [CH3CHO/Cu]+ at different O−X bond lengths.

Generally, the B3LYP predicts linear geometry for the CO− Al (173−180°), while bent arrangements (152−180°) were obtained by the MP2 calculations. There is no experimental structural data for [CH2O/Al]+ complex to assess the accuracy of the methods. However, the calculated gas-phase basicity values reported for formaldehyde toward Al+ (AlCB) at the B3LYP (89−93 kJ/mol) level are in better agreement with experimental AlCB (93.6 kJ/mol).5 The MP2-calculated AlCBs for CH2O using different basis sets are ∼79−87 kJ/mol. On the basis of the cation basicity of CH2O, the results of the B3LYP method (linear geometry) seem to be more accurate. Also,

comparable, and the CO−Al angle may be bent or linear depending on the computational method. The computational methods that predict a covalent nature for the O−Al interaction give a bent CO−Al angle, while linear geometry is obtained by methods that predict more electrostatic nature for the O−Al interaction. Therefore, on the one hand, the CO−Al angle in [CH2O/Al]+ depends on the computational methods and basis sets (Figure S2), so that it may be calculated from 152° to 180°. On the other hand, the CO−Cu and CO−Li angles do not depend on the computational method or basis sets and are, respectively, bent and linear at all levels of theory (Figure S2). D


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potential wells are located at ∼140° and 220° for O−Cu bond lengths of 1.9 (equilibrium bond length) and 3.5 Å (Figure 2d), indicating the dominance of the covalent contribution. At far enough distances, all ions are aligned with the dipole moment direction of the CH2O, regardless of the nature of the [CH2O/ X]+ interactions at the equilibrium geometry. As expected, at long distances the electrostatic contribution is the main part of the cation/molecule interaction. In the case of O−X bonds with covalent nature (O−Cu, e.g.), a longer O−X bond length is required for electrostatic contribution to become comparable with or larger than the covalent contribution. It should be mentioned that the cations and CH2O form planar structures, so that adduct ions with out-of-plane cations are not stable. Figure S3 shows the potential curves for rotation of Li+, K+, Al+, and Cu+ around CO axis. The results confirm that only the structure with planar >CO−X geometry are stable. The dipole moment axis of CH3CHO is slightly off from the direction of the CO bond (Figure 1), and the resulting C O−X angle scans exhibit small differences. Figure 3a shows the potential energy curves for CO−Li angle scan at O−Li bond lengths of 1.7, 2.5, and 5.0 Å for [CH3CHO/Li]+. At 1.7 Å, the potential well is located at ∼175°, where it is in the direction of the dipole moment axis of molecule. As the O−Li bond length increases (1.0−2.5 Å) the potential well shifts to smaller angles; then, at longer bond lengths (>2.5 Å) the minimum returns to larger angle, and finally, it remains constant at 172° (see also Figure S4). At 2.5 Å, the CO−Li angle of the most stable structure is ∼155°, which was not observed for formaldehyde (Figure 2a). The shift of the potential well with the O−Li bond length may be due to change in the electrostatic and covalent contributions as the O−Li bond length increases, which is consequently originated from change in the O−X distance and charge on X. Another possible interpretation can be proposed on the basis of the local or effective dipole moment that the cation feels at different distances from the molecule. Figure 3b shows the potential energy curves for [CH3CHO/K]+ at different O−K bond lengths. The shift in the potential well is not as much as that of [CH3CHO/Li]+, which may be due to electrostatic character of the K+ interactions. For [CH3CHO/ Al]+ and [CH3CHO/Cu]+ adduct ions, at equilibrium O−Al and O−Cu bond lengths and larger distances (up to 2.5 and 4.0 Å, respectively) the CO−Al and CO−Cu angles are smaller than 180° (Figure 3c,d). However, the CO−Cu angle shows more deviation from 180° implying a more covalent character of O−Cu bond relative to the O−Al bond. Furthermore, two minima are observed for [CH3CHO/Cu]+ at equilibrium O−Cu bond lengths and 140° and 220° indicating capability of Cu+ to interact covalently with both sides of CO group. However, the depths of the minima are not the same due to Cu/CH3 repulsion at CO−Cu angle of 220° (Figure 3d). The cation/CH3 repulsion presents in other [CH3CHO/ X]+ adduct ions and shows itself as asymmetry in potential curves at intermediate distances, which was not observed for [CH2O/X]+ complexes. At O−X distances longer than 5 Å, the cation/CH3 repulsion becomes negligible, and more symmetric curves are obtained. At far enough distances all cations are aligned with the dipole moment of CH3CHO to form [CH3CHO/X]+ adduct ions with CO−X angle of 172°. Acetone, (CH3)2CO, is the simplest ketone with dipole moment vector coinciding with the direction of the CO bond (similar to CH2O). Figure S5 shows the structures of [(CH3)2CO/X]+ adduct ions optimized by the B3LYP/6-311+

comparison of the B3LYP computed cation affinities and basicities with the experimental data (Table S1) shows the accuracy of the B3LYP method for the cation/molecule interactions. The differences between the computed and experimental data are smaller than ±10 kJ/mol, which are within the so-called chemical accuracy. For all cations (X) the O−X bond lengths in CH2O adduct ions are longer than those in CH3CHO adduct ions, in line with the stronger cation/ligand interaction in the latter case (Table S1). This may be attributed to the larger dipole moment of CH3CHO, which strengthens the electrostatic interaction, and an extra methyl group, which reinforces the covalent interaction.20 It should be mentioned that the methyl group increases both the electrostatic and covalent contributions. Addition of a methyl group increases the polarizability, which enhances the electrostatic ion/induced dipole interaction. Additionally, the electron-releasing effect of the methyl assists electron transfer toward the carbonyl and increases the covalent contribution. Table S3 summarizes the calculated ρ, ∇2ρ, G(r), and V(r) at the bond critical point (BCP) of the cation/molecule interactions. If we considered the electron density at BCP as a scale of covalency, it is concluded that the covalency of the [CH2O/X]+ and [CH3CHO/X]+ interactions depends on the cation type and varies as H+ > Cu+ > Al+ > Li+ > Na+ > K+. Furthermore, the −G(r)/V(r) ratios for alkali cations indicate that their interactions are mainly electrostatic. The negative sign of ∇2ρ for CH2O−H+ and CH3CHO−H+ shows that the covalent contributions of these interactions prevail. The positive values of ∇2ρ and negative sign of H(r) show partial covalency of the Cu+ and Al+ interactions. The results of AIM calculations confirm the relationship between the CO−X angle and nature of the cation/molecule interaction. To investigate in depth the relation between the nature of cation/molecule interaction and the structures of adduct ions, the CO−X angles in CH2O and CH3CHO adduct ions were scanned for different O−X bond lengths. Figures 2 and 3 show the potential energy curves obtained for CH2O and CH3CHO adduct ions, respectively. Figure 2a shows the curves obtained by scanning the CO−Li angle of [CH2O/Li]+ at O−Li bond lengths of 1.8, 2.5, and 4.0 Å. For O−Li bond length of 1.8 Å, the equilibrium bond length, a minimum, is observed at C O−Li angle of 180°, which is the global minimum of the adduct ion. As the bond length increases (2.5 Å) the potential well is broadened from 155° to 205°, with a minimum still at 180°. This broadening may be due to increase in the covalent contribution, which becomes comparable with the electrostatic interaction at 2.5 Å. However, the electrostatic contribution remains prevailing, and the energy minimum is observed in 180°. At far enough distances (4.0 Å) Li+ is aligned with the CO bond and dipole moment axis, so that the minimum appears at 180°. Figure 2b shows the potential energy curves obtained by scanning the CO−K angle of [CH2O/K]+ at O−K bond lengths of 2.5, 3.5, and 5.0 Å. At all bond lengths the potential wells are located at 180°, indicating the dominance of the electrostatic interaction at all distances. The curves of Figure 2c show that the potential well is displaced as the O−Al bond length increases. At O−Al bond length of 2.5 Å, the minima appear at 165° and 195°. This result shows that the ability of Al+ to participate in covalent interaction is more than that of Li+ and K+. At O−Al bond length of 5.5 Å, the [CH2O/Al]+ adduct ion with CO−Al angle of 180° is the most stable structure. For CH2O−Cu+, the E


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Figure 4. Optimized structures of Im−X+ adduct ions (X = Li+, Na+, K+, Al+, Cu+). The bond lengths (Å) and angles (deg) correspond to most stable structures optimized at the B3LYP/6-311++G(d,p) level of theory. The energies are in kilojoules per mole, and vector shows the direction of dipole moment of neutral imidazole-2-one.

Li+, Na+, and K+ show similar behaviors, and two isomers were obtained for each Im/alkali cation adduct ion. In the [Im/Li]+a, [Im/Na]+-a, and [Im/K]+-a the CO−X angles are 180°, and the cations are aligned with the direction of the molecule dipole moment. In the [Im/Li]+-b, [Im/Na]+-b, and [Im/K]+-b isomers, the cations are chelated by O and N atoms. Since the [Im/Li]+, [Im/Na]+, and [Im/K]+ interactions are mostly electrostatic and the axis of imidazolone dipole moment coincides with the direction of the CO bond, it is expected that isomers a are the most stable. Actually, the alkali cations interact with an effective dipole moment, resulting from local dipole moments induced by CO and CN groups, to form more stable isomers b (see relative energies in Figure 4). Since the covalent contribution of the Al+/molecule interactions is significant, Al+ can interact with both dipole moment of imidazolone ([Im/Al]+-a) and lone-pair electrons of N atom ([Im/Al]+-c). The CO−Al angle in the [Im/Al]+-b is ∼147°. This angle may be attributed to predominance of the covalent contribution; however, since the O−Al bond length in the [Im/ Al]+-b is slightly longer than that in the [Im/Al]+-a, probably the bending is largely due to local dipole moment of the adjacent imine group rather than covalent effect. However, both the electron density ρ and −G(r)/V(r) ratio at the BCP of Im− Al+ interaction show that the covalency of the interaction in the [Im/Al]+-b is larger than that for [Im/Al]+-a (Table S3). Although the bending of CO−Al in the [Im/Al]+-b may be due in part to the local dipole of the imine group, the bending could enhance the overlapping of the Al+ empty orbital with lone-pair electron of the atom O and increase the covalency of

+G(d,p) method. Despite CH3/cation repulsion, the covalent nature of [(CH3)2CO/H]+ and [(CH3)2CO/Cu]+ interactions result in bent CO−H and CO−Cu angles. For alkali cation complexes, linear CO−X geometries are formed due to the electrostatic character of the interactions. Since the covalent contribution to [(CH3)2CO/Al]+ is not significant, the CH3/Al+ repulsion prevails and leads in linear CO−Al geometry. Figure S6 compares the charge distribution and geometrical parameters of [(CH3)2CO/Cu]+ and [CH2O/ Cu]+. At equilibrium O−Cu bond length, the repulsion effect of Cu/CH3 is obviously larger than that of Cu/H; therefore, we expect a larger CO−Cu angle in [(CH3)2CO/Cu]+. The charge distribution indicates that the O−Cu interaction in [(CH3)2CO/Cu]+ is more covalent than that in [CH2O/Cu]+ due to larger polarizability of acetone and its extra electron donating methyl groups. Therefore, CO−Cu bending due to covalency overrides the Cu/CH3 repulsion, and the CO−Cu angle in [(CH3)2CO/Cu]+ (139°) is smaller than that in [CH2O/Cu]+ (143°). Although at short O−Cu distance, the Cu/CH3 interaction is mainly repulsion, Figure 3c shows that at longer O−Cu distance (3.5 Å), there is an attraction between Cu+ and CH3 group of acetaldehyde. 3.2. Imidazol-2-one (Imidazolone). Formaldehyde and acetaldehyde are two simple molecules with only one site to interact with the cations. Imidazolone (Im) contains one CO group and two N atoms, which are capable of interacting with cations and form chelate structures. Figure 4 shows the optimized structures for [Im/X]+ adduct ions (X+ = Li+, Na+, K+, Al+, Cu+) as well as their relative energies. The alkali cations F


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Figure 5. Potential energy curves obtained by scanning the CO−X angle of (a) [Im/Li]+, (b) [Im/K]+, (c) [Im/Al]+, and (d) [Im/Cu]+ at different O−X bond lengths.

the Im−Al+ interaction. In [Im/Al]+-c, only one BCP between Al and N was found; therefore, Al+ is not truly chelated by the N and O atoms, because, when it is located between O and N atoms, its covalent interaction with the N atom pulls it toward this atom ([Im/Al]+-c). Cu+ does not form chelate structure with imidazolone, and its interaction occurs only via N or O atom, [Im/Cu]+-a, and Im/Cu]+-b. Generally, nitrogen atoms tend to form more covalent bonds, while the electrostatic contribution to the CO−X+ interaction prevails as compared to that for N−X+ interactions. For these reasons, proton affinities of amines and imines are larger than those of

aldehydes and ketones, while alkali cation affinities of ketones are larger than those of amines and imines.5,25 Selective interaction of Cu+ with one of the N atoms of imidazolone (compare relative energies) indicates that the [Im/Cu]+ interaction is mainly covalent. The calculated electron densities ρ and −G(r)/V(r) ratios at the critical point of the [Im/X]+ interactions (Table S3) can be also used to compare the nature of the Im/X+ interactions. The ρ values for the Im/alkali cation interactions are smaller than those for [Im/H]+, [Im/Cu]+, and [Im/Al]+. Also, the −G(r)/V(r) ratios for Im−alkali cation interactions are greater than 1 indicating that they are mainly G


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Figure 6. Potential energy curves obtained by scanning the CO−X angle of (a) [γ-BL/Li]+, (b) [γ/BL-K]+, (c) [γ-BL/Al]+, and (d) [γ-BL/Cu]+ at different O−X bond lengths.

the adduct ions of imidazolin-2-one. The repulsion between the cations and H atom of N−H group results in only one isomer for each [imidazolin-2-one/X]+ complex. Linear geometries were observed for CO−Li, CO−Na, and CO−K angles due to electrostatic interaction between the cations and the dipole moment of the molecule, while bent CO−H and C O−Cu angles are because of overlapping of empty orbital of the cations and sp2 electron pairs of oxygen atom of CO group. As mentioned, the covalency of Al+ interactions is not as much as those of H+ and Cu+; therefore, H/Al repulsion overrides the

electrostatic. The cation affinity trend of imidazolone toward the studied cations is H+ > Cu+ > Li+ > Al+ > Na+ > K+ (Table S1). This trend is in accordance with the covalent nature of the interaction, except for Li+ and Al+. Although the covalency of the Al+ interactions is larger than those of Li+ interactions, the larger electrostatic contribution to the Li+ interaction results in higher LiCA compared to AlCA. Imidazolin-2-one is the saturated form of imidazole-2-one having two N−H groups instead of N atoms of imine groups in imidazole-2-one. Figure S5 shows the optimized structures of H


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Figure 7. Optimized structures of cytosine adduct ions of Li+, Na+, K+, Al+, and Cu+. The energies, bond lengths, and angles are in kilojoules per mole, angstroms, and degrees, respectively. The vector shows the direction of dipole moment of neutral cytosine.

O−X angle at long distances indicates that the electrostatic cation/dipole interaction becomes prevailing. 3.3. γ-Butyrolactone. Figure S7 shows the optimized structures of γ-butyrolactone (γ-BL) adduct ions of H+, Li+, Na+, K+, Al+, and Cu+. The calculated cation affinities of γ-BL were tabulated in Table S1. Esseffar et al.30 showed that γ-BL/ Cu+ complex has another isomer (interaction of Cu+ with CH2 group) with less stability than those shown in Figure S7. No chelate structure was observed for the γ-BL/cation adduct ions, so that the cations interact only with the oxygen atom of the CO group. Figure 6a shows the potential energy curves obtained by scanning the CO−Li angle of [γ-BL/Li]+ at different O−Li bond lengths. For O−Li bond length of 1.74 Å, the equilibrium bond length, a minimum, is observed at C O−Li angle of 162°, approximately the direction of the molecule dipole moment. As seen in Figure S7, at equilibrium O−X bond lengths, the alkali cations are not exactly aligned with the direction of the dipole vector of the molecule, and the CO−X angle is largely influenced by the local dipole of C O group. As the bond length increases (2.1 Å) the potential well shifts to smaller angles (135°) due to an increase of the covalent contribution. Effect of local dipole of the ether oxygen appears as a minimum at 90° (chelate structure). However, the chelate structure of [γ-BL/Li]+ is not a stable isomer at equilibrium bond lengths. At long O−Li distances (5.0 Å), Li+ is influenced by the total dipole moment of the molecule and aligns with its direction (146°). For the [γ-BL/K]+ complex, at O−K length of 2.5 Å the minimum is located at CO−K angle of 161° (Figure 6b). At longer O−K distances (2.8 Å) the minimum shifts to smaller angles (145°) due to increase in the covalency contribution. Also, for 2.8 Å a dip (not a true potential well) is observed at CO−K angle of 110°, which is

partial covalency, and an adduct ion with linear CO−Al angle is obtained. Figure 5 shows the potential energy curves obtained by scanning the CO−X angle of [Im/Li]+, [Im/K]+, [Im/Al]+, and [Im/Cu]+ at different O−X bond lengths. For [Im/Li]+ and [Im/K]+, at equilibrium bond lengths (1.78 and 2.5 Å, respectively), the minima of the potential wells are located at 180°. At O−Li bond length of 1.9 Å a minimum appears for [Im/Li]+ at CO−Li angle of 155° (Figure 5a), which may be attributed to a contribution of the covalent component rather than effect of the local dipole of the imine group, because bending of the CO−X angle due to the local dipole of the imine group results in chelate structures (the minima at ∼100° in Figure 5a,b). Figure 5c shows that, for [Im/Al]+ at O−Al bond length of 2.0 Å, the potential well is very flat between 160 and 200°, although three small depressions can be seen by enlarging this part of the curve. However, at O−Al length of 2.06 Å the potential wells at 150° and 210° become deeper and are related to the [Im/Al]+-b isomer (Figure 4). Since Cu+ has a strong covalent interaction with the N and O atoms ([Im/ Cu]+-a and [Im/Cu]+-b), the [Im/Cu]+ adduct ion with C O−Cu angle of 180° is not stable, and a maximum is observed for this geometry. At 1.94 Å, a minimum is observed at CO− Cu angle of 120° corresponding with the isomers [Im/Cu]+-b. At longer O−Cu distances a potential well is observed at angles smaller than 90° for the isomer [Im/Cu]+-a. Since the isomer [Im/Cu]+-a is more stable than the isomer [Im/Cu]+-b, the potential well at CO−Cu angles less than 90° is located at lower energies. For all adduct ions, except [Im/Cu]+, the structures with linear CO−X bond arrangement reach some stability at O−X bond length of 5.0 Å, while [Im/Cu]+ requires longer O−Cu bond length to reach this condition. Linear C I


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Figure 8. Optimized structures of adduct ions of CH2O, CH2S, and CH2Se with K+, Na+, Li+, Al+, Cu+, and H+ cations. Relative energies (when two stable forms are observed), bond lengths, and angles are in kilojoules per mole, angstroms, and degrees, respectively.

are related to two isomers of [γ-BL/Cu]+ complex (Figure S7). At O−Cu distance of 2.3 Å, a broad minimum is observed in the range of 100−120°. This broad minimum may result from merging of two minima, one for bending of CO−Cu angle attributable to the local dipole of the ether oxygen (a chelate structure) and another from bending of CO−Cu due to a dominance of covalent interaction. In the case of Li+ and K+ adducts, two distinguishable minima were observed instead of a broad potential well, because the covalent nature of O−Li and O−K bonds is weaker than that of the O−Cu bond, and

attributed to the local dipole of the ether oxygen. Appearance of the dip, rather than a minimum, may be due to larger size of K+ relative to Li+. In the case of [γ-BL/Al]+, at intermediate O− Al length (2.3 Å), only one minimum is observed at CO−K angle of 137° due to covalency contribution dominance (Figure 6c). Disappearance of the second minimum may be attributed to the deep minimum at 137° relative to that of chelate structure at smaller angles. For the [γ-BL/Cu]+ complex (Figure 6d), two minima (130 and 220°) are observed at O− Cu bond length of 1.9 Å, the equilibrium bond length, which J


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and orbital overlapping is enhanced. In fact, a cooperative effect between the better overlap at ∼140° and the orientation of the dipole moment may be advocated. This phenomenon is not observed for K+ and Na+, because K+ and Na+ essentially tend to undergo electrostatic interactions. In contrast, for Li+ and Al+, the interaction character depends on the structure of the host molecule. Formation of chelate structures depends on the cation type and the cation affinity of coordination centers. On the one hand, because of the covalent nature of Cu+ interactions, Cu+ interacts selectively with the nitrogen atom of imidazolone (the [Im/Cu]+-b isomer with O−Cu interaction is ∼35 kJ/mol less stable). On the other hand, the carbonyl group of cytosine is capable to participate in covalent interactions ([Cyt/Li]+-b; [Im/Al]+, [Cyt/Cu]+-b) similar to N atom; therefore, Cu+ forms a chelate structure with cytosine via the O and N atoms of carbonyl and imine groups. Therefore, chelate structures may be formed due to both covalent and electrostatic interactions. 3.5. CH2S and CH2Se. Figure 8 compares the optimized structures of adduct ions of CH2O, CH2S, and CH2Se with H+, Li+, Na+, K+, Al+, and Cu+ cations. Also, the dipole vectors were shown for CH2O, CH2S, and CH2Se. For the [CH2O/X]+ adduct ions, the CO−X angle is 180°, except for CO−Cu and CO−H. For [CH2S/X]+ adduct ions two energy minima (stable structures) with linear and bent CS−X geometries were obtained, the bent CS−X isomer being more stable. The relative energies of the two species are given for these isomeric forms. In the case of [CH2Se/X]+ adducts, only structures with bent CSe−X angle were obtained. The Mulliken charge distributions for [CH2O/X]+, [CH2S/X]+, and [CH2Se/X]+ are compared in Figure S10. The charge analysis results show that the [CH2S/X]+ and [CH2Se/ X]+ interactions are more covalent than the CH 2O/X + interaction. The trend is parallel to the polarizability order of Se > S > O. Since molecules with larger polarizability tend to form covalent cation/molecule bond, it is concluded that linear geometries are related to electrostatic charge/dipole interaction in CH2O/X+, and bent geometries indicate strong covalent interactions in the [CH2S/X]+ and [CH2Se/X]+. Comparison of the −G(r)/V(r) ratios in Table S2 indicates covalent nature of the [CH2S/X]+ and [CH2Se/X]+ interactions. Furthermore, Cu+ and H+ affinities of CH2S and CH2Se are larger than the corresponding values for CH2O, while for other cations the reverse result is observed (Table S1). Since the H+ and Cu+ mainly interact covalently, and dipole moments of CH2S (1.80 D) and CH2Se (1.54 D) are smaller than that of CH2O (2.46 D), it is concluded that the covalency of [CH2S/X]+ and CH2Se/X+ interactions are larger than that of [CH2O/X]+ interactions. In the case of CH2S and CH2Se, the CS− X and CSe−X angles are smaller than 120°; that is, they shift toward the angles observed in thiols, R−S−H, and in thioethers, R−S−R, which are closer to sp3 than to sp2 hybridization. Ammal and Venuvanalingum37,38 studied interaction of LiF with CH2Y (Y = O, S, Se) and reported the same trend for the C−Y−Li angles. They mentioned that the CH2O−LiF interaction is more electrostatic than the other two and showed that oxygen and sulfur donate nσ and nπ pairs, respectively. They attributed the smaller angle of C−S−Li to the nπ pairs donating (hybridization changing from sp2 to sp3) and hydrogen bonding interaction between F and hydrogen atoms of CH2 group.

therefore, C−O−Li and C−O−K angles are not considerably bent due to covalency. For all complexes, the CO−X angle approaches ∼147°, the direction of the γ-BL dipole moment (Figure S7). 3.4. Cytosine. Cytosine is one of the four important nucleobases in RNA and DNA. Its structure is more complex as compared to the simple carbonyls and imidazolone, bearing a CO group surrounded by three different N atoms (amine and imine groups). Cytosine has different keto and enolic isomers,31 and two isomers of them (a keto and an enol isomer) have considerable stability; however, only the most stable keto isomer is considered in this work. Figure 7 shows the optimized structures of adduct ions of cytosine, [Cyt/X]+, with X = Li, Na, K, Al, Cu. The calculated cation affinities of cytosine were compared with reported experimental and theoretical data in Table S1. There is a discrepancy between some of the reported experimental cation affinities and our calculated data. This discrepancy is attributed to different isomers of cytosine, keto and enolic forms.2 To clarify the origin of this discrepancy, the Li+, Na+, and K+ affinities of five stable isomers of cytosine were calculated (Figure S6). These calculations show that cation affinities of oxygen atom of cytosine are larger in the keto isomers relative to the enolic forms. Therefore, the smaller values of cation affinities are related to the enolic isomer. The alkali cations Li+, Na+, and K+ are chelated in the cytosine adducts by the O atom of carbonyl group and the adjacent N atom of the imine group, while Al+ interacts only with the O atom of CO group. These structures have been previously reported for the cytosine adduct ions.18,19,32−36 Cu+ is chelated by the O and N atoms ([Cyt/Cu]+-a); however, a comparison of the N−Cu and O− Cu bond lengths with N−X and O−X bonds in the chelate structures of other cations shows that Cu+ favors N bonding, in contrast to the alkali cations, which favor O bonding. In other words, because of covalent character of bonding to Cu+ the interaction tends to be stronger with the nitrogen site (see ρ and −G(r)/V(r) values in Table S3). Also, Cu+ interacts with oxygen atom of CO group of cytosine to form another isomer ([Cyt/Cu]+-b), which is less stable than [Cyt/Cu]+-a by ∼5.8 kJ/mol. The CO−Cu angle of the former isomer (∼115.5°) indicates orbital overlapping and strong covalent character of the O−Cu interaction. Since no BCP was found between Li and the N atom in the [Cyt/Li]+-b, with a CO−Li angle of ∼136°, this isomer cannot be considered as a true chelate structure. This structure is similar to the [Cyt/Al]+ adduct ion with a CO−Al angle of ∼146°. The bending of the CO−Li and CO−Al may be attributed in part to the covalent nature of the O−Li and O−Al bonds, but the strong dipole moment of cytosine (6.7 D), which passes through the Li and Al atoms in [Cyt/Li]+-b and [Cyt/Al]+, should also play a role. The considerable values of ρ at the BCP of the O−Al and O−Li bonds (Table S3) and Mulliken charge distribution analysis (Figure S9) shows that the O−Li and O−Al bonds in the [Cyt/Li]+-b and [Cyt/Al]+ have a sizable covalent character (partial charges on Li and Al are 0.65 and 0.40, respectively). Furthermore, comparison of the partial charges on the Li and Al atoms in the formaldehyde and cytosine adducts (Figures S1 and S9) reveals that the covalent character of the O−Li and O−Al bonds in cytosine adducts is larger. One reason is related to the larger basicity of cytosine compared to formaldehyde. In the [Cyt/Li]+-b and [Cyt-/Al]+, the CO−Li and CO−Al angles are bent by local dipole of the imine group or total dipole of the cytosine, K



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4. CONCLUSION The Lewis acids Cu+ and H+ interact mainly covalently with one of the electron lone pairs of the carbonyl group (sp2); hence, they form an adduct with a bent CO−X geometry. The alkali metal cations tend to participate in electrostatic interactions; therefore, they tend to align with the direction of dipole moment of carbonyl group with CO−X angle close to 180°. However, the CO−X angle may deviate from linearity due to substitution or local dipole moment of an adjacent group. The electrostatic and covalent contributions of Al+ interactions are both significant; inducing a CO−X bending results in a better orbital overlap and pronounced covalency, like in [Im/Al]+-b and [Cyt/Al]+. Even for Li+ whose interactions are seen as largely electrostatic, bending of C O−Li angle increases the covalency of the CO−Li interaction, as in [Cyt/Li]+-b. In addition, we showed that the nature of CO−X+ interaction varies with the O−X bond lengths, which consequently influences the CO−X angle. Generally, at intermediate O−X distances, the covalent contribution prevails over the electrostatic and the bent CO−X geometry becomes more stable. At long distances the interaction becomes mainly electrostatic, as expected, and the cation aligns with the direction of the dipole moment of the molecules. The importance of this result is in chelate structures, large molecules such as proteins with several acceptor sites, and supramolecular assemblages. These structures undergo deformations from the initial single ligand/cation interaction. From a given cation attached to a given site and considering the distances of other surrounding sites, we may estimate what is the direction of the preferred distortion. Also, the nature of interaction (covalent or electrostatic) between the cation and surrounding sites can be predicted based on the distance of the cation from that site.

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ASSOCIATED CONTENT

S Supporting Information *

The Supporting Information is available free of charge on the ACS Publications website at DOI: 10.1021/acs.jpca.7b04474. Cation affinities, ionization energies of the metals, charge distribution of formaldehyde, effect of computational method on CO−Al angle, the calculated ρ, ∇2ρ, G(r), and V(r), potential energies of [CH2O/X]+, structures of [CH3CHO/X]+ adduct ions, structures of adduct ions of acetone and imidazolin-2-one, charge distribution of [(CH3)2CO/Cu]+ and [CH2O/Cu]+, structures of γbutyrolactone adduct ions, structures of alkali metal cation adducts of cytosine, charge distribution for the cytosine adduct ions, and charge distribution for the CH2O, CH2S, and CH2Se adduct ions (PDF)

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AUTHOR INFORMATION

Corresponding Author

*E-mail: [email protected]. Phone: +98 21 4486 5001. ORCID

Younes Valadbeigi: 0000-0002-4189-2987 Notes

The authors declare no competing financial interest.

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ACKNOWLEDGMENTS J.F.G. thanks the French National Agency for Research (Grant No. ANR ALEA-2013-009-01) for partial support of this work. L


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