CarbonâCarbon Bonding between Nitrogen Heterocyclic Carbenes

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Article Cite This: J. Phys. Chem. A XXXX, XXX, XXX-XXX

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Carbon−Carbon Bonding between Nitrogen Heterocyclic Carbenes and CO2 Published as part of The Journal of Physical Chemistry virtual special issue “W. Lester S. Andrews Festschrift”. Janet E. Del Bene,*,‡ Ibon Alkorta,*,§ and José Elguero§ ‡

Department of Chemistry, Youngstown State University, Youngstown, Ohio 44555, United States Instituto de Química Médica, CSIC, Juan de la Cierva, 3, E-28006 Madrid, Spain

§

S Supporting Information *

ABSTRACT: Ab initio MP2/aug′-cc-pVTZ calculations were performed to identify equilibrium complexes and molecules and the transition structures that interconvert them, on the potential energy surfaces of a series of seven binary systems that have nitrogen heterocyclic carbenes (NHCs) as the electron-pair donors to CO2. Seven of the NHCs form complexes stabilized by C···C tetrel bonds, and six of these seven are also stabilized by a secondary interaction between an O of CO2 and the adjacent N−H group of the carbene. Six of the seven NHCs also form stable molecules with C−C covalent bonds, and with one exception, these molecules have binding energies that are significantly greater than the binding energies of the complexes. Charge-transfer stabilizes all of the NHC:CO2 complexes and occurs from the C lone pair of the carbene to the CO2 molecule. The six complexes that have secondary stabilizing interactions are also stabilized by back-donation of charge from the O to the adjacent N−H group of the carbene. Transition structures present barriers to the interconversion of complexes and molecules. With one exception, the barrier for converting a molecule to a complex is much greater than the barrier for the reverse reaction. Atoms in Molecules bonding parameters, shifts of IR C−O stretching and O−C−O bending frequencies, changes in NMR 13C chemical shieldings, and changes in C−C and C−O coupling constants as 1tJ(C−C) and J(C− O) for complexes and transition structures become 1J(C−C) and 2J(C−O) for molecules, are all consistent with the changing nature of the C···C tetrel bond in the complex through the transition state to a covalent C−C bond in the molecule.

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INTRODUCTION

There has been a resurgence of interest in the tetrel bond, as evidence by the number of papers published recently on this subject. In 2014, Frontera et al. noted that there is a σ-hole between the carbon atoms in cycloalkane structures such as (CN)2C−C(CN)2, so that electron-rich donors can form stable complexes with these molecules.21 Later, Guo et al. examined competition and cooperativity involving tetrel bonds and chalcogen bonds.22 Scheiner reported a study of CH···O, SH···O, chalcogen, and tetrel bonds formed by neutral and cationic sulfur-containing compounds.23 Two studies of cooperativity by Esrafili et. al and Marin-Luna et al. in chains of acetonitrile and isoacetonitrile and the analogous silyl derivatives were also reported.24,25 Subsequently, Bauzá and Frontera examined bonding through C−C σ-holes and πholes.26 The role of the tetrel bond in SN2 reactions has been studied by Liu et al.27 We reported studies of hydrogen and halogen bonds between anions and substituted methanes.28 In 2017, studies of carbon−carbon interactions involving carbenes as electron donor were reported by Liu et al.29 In addition,

The field of weak interactions is continuously expanding to include additional elements of the periodic table. Hydrogen bonding has traditionally dominated this field and will probably continue to do so simply because of the importance of hydrogen bonding in water and biological systems. The weak interaction that is probably next in terms of familiarity is the halogen bond, which was discussed in detail in the pioneering work of Legon,1,2 and has been the subject of recent review articles.3−5 This bond arises when a group 17 halogen atom acts as an electron-pair acceptor through its σ-hole.6,7 Other types of σ-hole interactions exist, including chalcogen,8−11 pnicogen,12−14 and tetrel bonds.15−18 These bonds arise when elements from groups 16, 15, and 14, respectively, act as electron-pair acceptors. Legon and Resnati et al. have recently suggested that all of these bonds should be considered as arising when a σ-hole or a π-hole associated with an E atom in one molecular entity interacts with a nucleophilic region such as a pair of nonbonding or π electrons in another, or the same, molecular entity.19 It is interesting to note that the discovery of these interactions occurred earlier than their naming, as noted in a recent review by Legon.20 © XXXX American Chemical Society

Received: August 22, 2017 Revised: September 22, 2017

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DOI: 10.1021/acs.jpca.7b08393 J. Phys. Chem. A XXXX, XXX, XXX−XXX

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The Journal of Physical Chemistry A

Dunning aug-cc-pVTZ basis set40,41 by removing diffuse functions from H atoms. Frequencies were computed to establish that the structures of complexes and molecules are equilibrium structures with no imaginary frequencies and that the transition structures have one imaginary frequency along the coordinate which interconverts the two equilibrium structures. Optimization and frequency calculations were performed using the Gaussian 09 program.42 Intrinsic reaction coordinate (IRC) calculations43,44 were performed to illustrate the path connecting these structures for the 1:CO2, 4:CO2, and 7:CO2 systems. The binding energies of the binary NHC:CO2 complexes and NHC−CO2 molecules were evaluated as −ΔE for the reaction that forms the complex or the molecule from the corresponding NHC and CO2 monomers. For consistency, the binding energies of transition structures were evaluated in the same way, which means that a positive binding energy indicates that the transition structure is bound relative to the corresponding isolated NHC and CO2. The electron densities of complexes and molecules were analyzed using the Atoms in Molecules (AIM) methodology45−48 employing the AIMAll49 program. The topological analysis of the electron density produces the molecular graph of each. This graph identifies the location of electron density features of interest, including the electron density (ρ) maxima associated with the various nuclei, and saddle points that correspond to bond critical points (BCPs). The zero gradient line that connects a BCP with two nuclei is the bond path. The electron density at the carbon−carbon bond critical point (ρBCP), the Laplacian (∇2ρBCP) at that point, and the total energy density (HBCP) were also computed. The molecular electrostatic potentials (MEPs) were also evaluated for the seven NHCs with the DAMQT program.50 The Natural Bond Orbital (NBO) method51 has been used to obtain the stabilizing charge-transfer interactions for complexes using the NBO 6 program.52 Since MP2 orbitals are nonexistent, charge-transfer interactions were computed using the B3LYP functional with the aug′-cc-pVTZ basis set at the MP2/aug′-cc-pVTZ complex geometries. This allows for the inclusion of at least some electron correlation effects. NMR absolute chemical shieldings were evaluated at the MP2/aug′-cc-pVTZ level with the Gauge-Invariant Atomic Orbital (GIAO) method53,54 using the Gaussian 09 program. Equation of motion coupled cluster singles and doubles (EOMCCSD) spin−spin coupling constants were evaluated in the configuration interaction (CI)-like approximation55,56 with all electrons correlated. For these calculations, the Ahlrichs57 qzp basis set was placed on 13C, 15N, and 17O, and the Dunning ccpVDZ basis set40 was placed on 1H atoms. Total coupling constants were evaluated as the sum of the paramagnetic spin orbit (PSO), diamagnetic spin orbit (DSO), Fermi contact (FC), and spin dipole (SD) terms using ACES II58 on the HPC cluster Oakley at the Ohio Supercomputer Center.

several papers by Scheiner and his group appeared at this time in which several different aspects of tetrel bonding were addressed. These included the factors that influence the strengths of tetrel bonds,30 a comparison of hydrogen, halogen, chalcogen, and tetrel bonds,31 and a comparison of tetrel bonds in neutral complexes of pyridine and furan with all of the group 14 elements.32 We also published two papers in 2017 on reactions involving carbenes and CO2.33,34 In the first, a variety of singlet carbene bases including CF2, CCl2, CCH2, CCCH2, cyclic C3H2, C(CH3)2, C(OH)2, C(NH2)2, C(OCH)2, and C(NHCH)2 were electron-pair donors to CO2 for the formation of C···C tetrel bonds and C−C covalent bonds. The second article focused on the interaction of CO2 with seven N,Ndimethylated nitrogen heterocyclic carbenes (NHCs) and on the relationship between the binding energies of these complexes and the orbitals and molecular electrostatic potentials (MEPs) of the isolated carbene bases. In the present article we examine the structures and binding energies of a series of seven related NHC:CO2 systems involving the parent NHC molecules. The seven NHC bases are illustrated in Scheme 1. NHCs 1, 2, and 3 are referred to as Scheme 1. Seven Nitrogen Heterocyclic Carbenes

Classical NHCs; 4, 5, and 6 are Abnormal NHCs; and 7 is a Remote NHC. We also report bonding properties and spectroscopic properties of the complexes, molecules, and transition structures found on the NHC:CO2 potential surfaces. The spectroscopic properties include selected IR frequencies and NMR chemical shieldings and spin−spin coupling constants.

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METHODS The seven nitrogen heterocyclic carbenes, the CO2 molecule, and the NHC:CO2 complexes, NHC−CO2 molecules, and NHC/CO2 transition structures were optimized at secondorder Møller−Plesset perturbation theory (MP2)35−38 with the aug′-cc-pVTZ basis set.39 This basis set was derived from the

Scheme 2. Reaction between NHC 1 and CO2 Illustrating the Formation of Complex, Transition Structure, and Molecule

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RESULTS AND DISCUSSION Scheme 2 illustrates the reaction involving NHC 1 and CO2 to form the 1:CO2 complex, the 1-CO2 molecule, and the 1/CO2 transition structure that interconverts them. The IRC reaction path connecting these structures is shown in Figure S1 of the Supporting Information. The first section below provides a discussion of the isolated carbenes. This is followed by a discussion of the structures, binding energies, and chargetransfer energies of the complexes in section two and of the structures and binding energies of molecules and transition structures in sections three and four, respectively. Section five focuses on the evolution of the electron density parameters at C4−C1 bond critical points as complexes pass through transition states to become molecules. In these systems, C1 is the carbon atom of CO2, and C4 is the carbene carbon that is the electron-pair donor to C1; the oxygen atoms are O2 and O3. The final section deals with spectroscopic properties, including IR stretching and bending frequencies of the CO2 molecule, changes in 13C1 and 13C4 NMR chemical shieldings along the reaction coordinate, and NMR C4−C1, C4−O2, and C4−O3 spin−spin coupling constants for the three stationary points of interest. NHC Monomers. The isolated NHCs 1, 3, and 7 have C2v symmetry, while 2, 4, 5 and 6 have Cs symmetry, as can be seen in Table S1 of the Supporting Information. Their MEPs in the singlet electronic state were evaluated, and the MEPs for NHCs 4 and 7 are illustrated in Figure 1. Each MEP has a minimum

is the O atom of CO2 that may interact with N5−H or C5−H of the carbene. C1 is the carbon atom of CO2. Table S2 of the Supporting Information presents the structures, total energies, and molecular graphs of the NHC:CO2 complexes, all of which have Cs symmetry. Table 2 reports their binding energies and intermolecular C4−C1 distances. The complex 4:CO2 has the highest binding energy of 29.6 kJ·mol−1 at the shortest intermolecular C4−C1 distance of 2.792 Å, while 3:CO2 has the lowest binding energy of 20.4 kJ·mol−1 at a C4−C1 distance of 3.001 Å. Nevertheless, the correlation between the binding energies and the C4−C1 distance is not good, as evident from Figure 3. From this plot it appears that the binding energy of 5:CO2 is too large for its C4−C1 distance, while that of 7:CO2 is too small. The correlation coefficient of the exponential trendline is 0.863. Figure 4 presents a plot of the binding energies versus the MEP minima of the NHC bases for these same complexes. A muchimproved second-order correlation with a correlation coefficient of 0.982 is found but only if the point for 7:CO2 is omitted. Why are good correlations for the entire set of complexes not found? Insight into the answer to this question may be found by considering the orientation of the CO2 molecule in these complexes. From the molecular graphs in Table S2, it is evident that the CO2 molecule is not perpendicular to the C4−C1 axis. Rather, CO2 is slightly bent and tilted to facilitate interaction between O2 and N5−H in complexes of CO2 with NHCs 1−6, and between O2 and C5−H in 7. In complexes 1−6, the H− N5−O2 angle lies between 42 and 44°, indicating that this interaction is not a distorted N5−H···O2 hydrogen bond but is better described as a stabilizing electrostatic interaction. In contrast, the H−C5−O2 angle in 7:CO2 is 57°, and it is a C−H group that is interacting with O2, which in itself reduces the electrostatic interaction. Moreover, the distances between O2 and the hydrogen atom bonded to N5 in complexes with NHCs 1−6 are between 2.21 and 2.31 Å, while the corresponding distance in 7:CO2 is 2.67 Å. These data suggest that the absence of this secondary interaction in 7:CO2 is responsible for its lower relative stability compared to the complexes with NHCs 4, 5, and 6, even though 7 has the most negative MEP value. It is of interest to compare the binding energies of the N,Ndimethylated NHC complexes with CO2 from ref 34 with those of the hydrogenated NHC:CO2 complexes investigated in this work. The binding energies of the dimethylated complexes are G4MP4 energies, and they are reported in Table S3 of the Supporting Information. While these energies are not strictly comparable to the MP2/aug′-cc-pVTZ binding energies of the parent complexes, the trends in the two series should be similar. Omitting complexes with 7 for the moment, the order of binding energies for the two series is similar, except for a reversal of complexes with 1 and 6. However, the binding energies of these two methylated NHC:CO2 complexes differ by only 0.2 kJ·mol−1, while those of the parent complexes differ by 0.5 kJ·mol−1. What is most interesting, however, is that the methylated 7:CO2 complex has the largest binding energy among all of the methylated complexes. This is in contrast to

Figure 1. MEP +0.10 au (blue) and −0.10 au (red) isosurfaces for NHCs 4 and 7.

associated with the lone pair of electrons on the carbene C4, and the values of these minima are given in Table 1. NHC 3 has the least negative MEP minima, most probably because it has four N atoms in the ring, which can withdraw electrons from C4. NHC 7 has the most negative value, since the N−H groups are not bonded to C4 and would therefore have the smallest electron-withdrawing effect on the C4 lone pair. Whether or not the MEP values correlate with the binding energies of the NHC:CO2 complexes will be discussed below. Structures, Binding Energies, and Charge-Transfer Energies of NHC:CO2 Complexes. All seven NHCs form binary complexes with CO2, and six of the seven form bound molecules as well. Figure 2 illustrates the complex, transition structure, and molecule formed by NHC 7 and CO2. In structures with Cs symmetry, O2 and O3 are not equivalent. O2 Table 1. MEP Minima Values (au) for the Seven NHCs NHC

1

2

3

4

5

6

7

MEP-min

−0.127

−0.110

−0.090

−0.148

−0.139

−0.126

−0.155

C


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Figure 2. Complex, transition structure, and molecule formed by NHC 7 and CO2. In systems with Cs symmetry, O2 is the O atom of CO2 that may interact with N5−H or C5−H of the NHC.

Table 2. Binding Energies (−ΔE) and Charge-Transfer Energies (CT, kJ·mol−1) and C4−C1 Distances (R, Å) for NHC:CO2 Complexes NHC

−ΔE

R(C4−C1)

CT (C4lp→π*C1−O3)

CT (O2lp→σ*N5−H)

1 2 3 4 5 6 7

24.8 22.7 20.4 29.6 28.6 25.3 25.0

2.876 2.939 3.001 2.792 2.848 2.897 2.838

17.0 12.3 8.8 30.0 23.5 17.5 29.3

2.3 2.4 2.1 3.3 3.4 2.6

between O2 and the N5−H group. The presence of two methyl groups eliminates this secondary interaction in the methylated complexes. Table 2 also provides the stabilizing C4lp→π*C1−O3 charge-transfer energies for the NHC:CO2 complexes, and Figure 5 illustrates the orbitals involved in charge transfer in three of these complexes. Charge-transfer energies vary from 9 kJ·mol−1 for the most weakly bound 3:CO2 complex to 30 kJ· mol−1 for the most strongly bound 4:CO2. The exponential dependence of the charge-transfer energies on the C4−C1 distance has a correlation coefficient of 0.951, and it is also illustrated in Figure 3. Charge transfer depends on the ability of the NHC to donate electrons, the ability of CO2 to accept an electron pair, and on the C4−C1 distance, with no other secondary interactions being significant. 7:CO2 has the secondlargest charge-transfer energy of 29 kJ·mol−1. This is consistent with the value of the MEP minimum for 7 and its short C4−C1 distance. Thus, the point for 7:CO2 is included in this plot. There is also a second charge-transfer interaction in complexes with NHCs 1 through 6. This interaction is a back-donation of charge from the O2 lone pair to the antibonding σ N5−H orbital of the carbene. This second charge-transfer interaction is consistent with the existence of a secondary interaction between O2 and N5−H in these complexes. The absence of back-donation of charge in 7:CO2 is further support that the reduced binding energy of this complex is not due to a weakened tetrel bond but to the absence of a secondary interaction that adds stability to the complexes of CO2 with NHCs 1 through 6. Structures and Binding Energies of NHC−CO 2 Molecules. Table S4 of the Supporting Information presents the structures, total energies, and molecular graphs of the NHC−CO2 molecules. All of the molecules are stable relative to the corresponding isolated monomers NHC and CO2, except for 3-CO2, which has a binding energy of −14.2 kJ· mol−1 and will not be considered further. Molecules 1-CO2 and 7-CO2 have C2v symmetry, while the remaining molecules have Cs symmetry. Molecules 1-CO2 and 6-CO2 are shown in Figure

Figure 3. Binding energies and charge-transfer energies of NHC:CO2 complexes vs the C4−C1 distance.

Figure 4. Binding energies of NHC:CO2 complexes vs the MEP minima of NHCs.

the hydrogenated complexes, in which case 4:CO2, 5:CO2, and 6:CO2 are more stable than 7:CO2. As noted above, the parent complexes are stabilized by a secondary electrostatic interaction D


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Figure 5. Orbitals involved in the C4lp→π*C1−O3 charge transfer for 1:CO2, 4:CO2, and 7:CO2.

6, and molecular binding energies and C4−C1 distances are reported in Table 3. These binding energies range from 23 to

Figure 6. Molecules 1-CO2 and 6-CO2.

83 kJ·mol−1, while the C4−C1 distances vary by less than 0.02 Å from 1.524 to 1.543 Å, values typical of C−C single covalent bonds. The largest molecular binding energy is not found at the shortest C4−C1 distance, nor is the smallest binding energy found at the longest distance, as clearly illustrated by the scattergram given as Figure 7. The lack of correlation between these two variables is a result of the differences in the bonding schemes in the NHC rings and therefore in their electron density distributions. The binding energies of the NHC−CO2 molecules are significantly greater than the binding energies of the corresponding NHC:CO2 complexes, except for NHC 2, in which case the complex and the molecule have the same binding energy. The binding energies of the N,N-dimethylated NHC−CO2 molecules are also reported in Table S3 of the Supporting Information. Once again, 7-CO2 is the most stable among these, whereas 4-CO2 and 5-CO2 are more stable than 7-CO2 among the parent molecules. The presence of the methyl groups also lengthens the C4−C1 bond in the N,Ndimethylated complexes relative to that bond in the parent NHC−CO2 molecules. Structures and Binding Energies of Transition Structures. Table S5 of the Supporting Information reports the structures, total energies, and molecular graphs of the NHC/CO2 transition structures. Table 4 presents their binding energies and C4−C1 distances. All of the transition structures except 2/CO2 have positive values of −ΔE; that is, they are bound relative to isolated NHC and CO2. Table 4 also presents the barriers to the forward and reverse reactions for interconverting the complex and the molecule. Except for 2CO2, the barrier for the molecule going to the complex is

Figure 7. Binding energies vs the C4−C1 distance for NHC−CO2 molecules.

always greater, since the molecule is found in a deeper minimum on the potential surface. The barriers for the forward and reverse reactions for NHC 2 are the same, since the binding energies of the complex and the molecule are the same. How does the structure change as the complex crosses through the transition structure to become a molecule? Tables 2, 3, and 4 and Tables S2, S4, and S5 can assist in answering this question. There are two structural changes that are immediately apparent, namely, the bending of the CO2 molecule and the change in the C4−C1 distance. On the 7CO2 surface, the O2−C1−O3 angle decreases from 174° in the complex to 157° in the transition structure and 134° in the molecule. This bending is evident from Figure 2, and it is observed on all potential surfaces. On the same surface, the C4−C1 distance decreases dramatically from 2.838 Å in the complex to 2.226 Å in the transition structure and 1.542 Å in the molecule. It is the change in the C4−C1 distance that can be directly related to the barrier height. This can be seen from two different distance relationships. The first is the parameter β,59 which is defined as δR1/δR2, with δR1 = R ts − R molecule δR 2 = R complex − R molecule

The β values in Table 4 that are less than 0.50 are associated with transition states that are weakly bound or even unbound, as is the case for NHCs 1, 2, and 6. Systems with larger β values have more strongly bound transition structures. Figure 8

Table 3. Binding Energies (−ΔE, kJ·mol−1) and C4−C1 Distances (R, Å) for NHC−CO2 Molecules NHC

1

2

4

5

6

7

−ΔE R(C4−C1)

57.2 1.524

22.8 1.535

82.7 1.534

73.3 1.539

40.4 1.543

70.4 1.542

E


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Table 4. Binding Energies (−ΔE) and Transition-State Barriers (E‡, kJ·mol−1), C4−C1 distances (R, Å) for NHC/CO2 Transition Structures, and C4−C1 Distance Relationships [β and ΔR, Å] among Complexes, Transition Structures, and Molecules NHC

−ΔE

R(C4−C1)

E‡a

E‡b

βc

ΔR(Rcomplex − Rts)d

ΔR|Rmolecule − Rts|d

1 2 4 5 6 7

9.3 −5.8 25.2 21.6 8.0 17.0

2.113 1.990 2.285 2.246 2.083 2.226

15.5 28.5 4.4 7.0 17.3 8.0

42.1 28.6 57.5 51.7 32.4 53.4

0.436 0.324 0.597 0.540 0.399 0.528

0.763 0.949 0.507 0.602 0.814 0.612

0.589 0.455 0.751 0.707 0.540 0.684

Barrier for the reaction complex → molecule. bBarrier for the reaction molecule → complex. cβ = δR1/δR2; δR1 = Rts − Rmolecule; δR2 = Rcomplex − Rmolecule. dThe distance in bold is closer to the transition structure distance. a

Table 5. Electron Densities (ρBCP), Laplacians (∇2ρBCP), and Total Energy Densities (HBCP) (au) at C4−C1 Bond Critical Points of Complexes, Transition Structures, and Molecules

Figure 8. Barriers for the forward (complex→molecule) and reverse (molecule→complex) reactions vs the parameter β.

illustrates the linear correlations that exist between the barriers for the forward and reverse reactions and the parameter β. The correlation coefficients are 0.960 and 0.963, respectively. The second distance relationship is the difference between the C4−C1 distance in the complex and the transition structure and the difference between the C4−C1 distance in the molecule and the transition structure. These differences are also reported in Table 4. When the C4−C1 distance in the transition structure is closer to the C4−C1 distance in the complex, the barrier for converting the complex to the molecule is low, with values between 4 and 8 kJ·mol−1 for NHCs 4, 5, and 7. The highest barriers between 52 and 58 kJ·mol−1 are those associated with the reverse reaction in these same systems. When the C4−C1 distance in the molecule is closer to that in the transition structure in systems with NHCs 1, 2, and 6, the barriers for converting the molecule to the complex are between 29 and 42 kJ·mol−1, smaller than those for the same reaction involving NHCs 4, 5, and 7. The barriers are between 16 and 29 kJ·mol−1 for converting the complex to the molecule in these same systems. AIM Bonding Data. C4−C1 bond critical point data including electron densities (ρBCP), Laplacians (∇2ρBCP), and total energy densities (HBCP) for complexes, transition structures, and molecules are reported in Table 5. Each of these bonding parameters shows the evolution of the C4−C1 bond as the complex crosses the transition state and becomes a molecule. This is illustrated by the graphs of Figure S2 in the Supporting Information. The electron densities at BCPs are less than 0.02 au in complexes, increase to ∼0.08 au in transition

complexes

ρBCP

∇2ρBCP

1:CO2 2:CO2 3:CO2 4:CO2 5:CO2 6:CO2 7:CO2 transition structures 1/CO2 2/CO2 4/CO2 5/CO2 6/CO2 7/CO2 molecules 1-CO2 2-CO2 4-CO2 5-CO2 6-CO2 7-CO2

0.0141 0.0123 0.0108 0.0173 0.0157 0.0139 0.0155

0.0409 0.0380 0.0356 0.0444 0.0416 0.0400 0.0412

0.0006 0.0009 0.0011 0.0000 0.0003 0.0006 0.0003

0.0685 0.0878 0.0488 0.0534 0.0732 0.0546

0.0582 0.0380 0.0610 0.0587 0.0510 0.0601

−0.0178 −0.0294 −0.0091 −0.0110 −0.0204 −0.0113

0.2479 0.2400 0.2436 0.2418 0.2367 0.2384

−0.6734 −0.6297 −0.6392 −0.6250 −0.6048 −0.6102

−0.2473 −0.2364 −0.2329 −0.2285 −0.2254 −0.2228

HBCP

structures, and then increase again in molecules to 0.25 au. Figure S2a shows the exponential relationship with a correlation coefficient of 0.997 for the variation of the electron densities at C4−C1 bond critical points as a function of the C4−C1 distance. The Laplacians are positive at long distances, increase slightly as the distance decreases, reach a maximum, and then decrease and become negative as the distance continues to decrease. The variation of the Laplacians as a function of the C4−C1 distance can be seen in Figure S2b. Finally, the energy densities are positive but have values of 0.001 au or less for the complexes, become negative with values of ca. −0.02 au for the transition structures, and then further decrease to between −0.22 and −0.25 au for the molecules. This behavior is illustrated in Figure S2c. These bonding data reflect the changes in the C4−C1 bond from an intermolecular tetrel bond at long distances in the complexes, to an intermolecular bond with some covalent character at the intermediate distances in the transition structures, to a covalent bond at the short distances found in the molecules.60,61 Spectroscopic Properties. IR Spectra. The CO2 molecule exhibits symmetric and asymmetric C−O bond stretches, and in-plane and out-of-plane O−C−O bends. In transition F


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The Journal of Physical Chemistry A Table 6. Values of the CO2 Asymmetric Stretching and In-Plane Bending Frequenciesa asymmetric stretching, cm−1

a

in-plane bending, cm−1

NHC

NHC:CO2

NHC/CO2

NHC−CO2

NHC:CO2

NHC/CO2

NHC−CO2

1 2 3 4 5 6 7

2396 2400 2403 2386 2391 2397 2389

2180 2114

1772 1792

542 572

383 367

2249 2229 2154 2216

1765 1764 1785 1760

613 624 634 587 597 612 592

519 525 558 510

385 385 371 374

The asymmetric stretching and in-plane bending frequencies of isolated CO2 are 2401 and 659 cm−1, respectively.

monomer decreases to between 587 and 634 cm−1 in the complexes, between 510 and 572 cm−1 in the transition structures, and between 367 and 385 cm−1 in the molecules. In the complexes, the out-of-plane bending frequency decreases as the C4−C1 distance decreases, and these two variables are linearly related with a correlation coefficient of 0.95. The same band in the transition structures increases linearly as the C4−C1 distance decreases, with a correlation coefficient of 0.94. Note, however, that this band is closely related to the imaginary frequency of the transition structures. This band in the molecules does not correlate with changes in the C4−C1 distance. 13 C NMR Chemical Shieldings. The chemical shieldings of the carbon atom C1 of the CO2 molecule and C4 of the carbene are reported in Table 7. The chemical shielding of 13C1 decreases relative to CO2 at the three stationary points on the surface. The decrease is very small at −0.6 to −0.9 ppm in the complexes, larger for the transition structures at −3.3 to −6.4 ppm, and is largest in molecules at −20.7 to −25.2 ppm. The change in the 13C1 chemical shieldings versus the C4−C1 distance is illustrated in Figure 10. The exponential relationship between these two parameters has a correlation coefficient of 0.98. In contrast, the chemical shieldings of 13C4 of the carbene increase upon complexation, except for the 7-CO2 molecule, where a deshielding of 1.5 ppm is found. Moreover, interaction of the carbene with CO2 has a greater effect on the chemical shieldings of C4 than C1, except for the stationary points involving NHC 7. Relative to the isolated NHCs, the chemical shieldings of C4 increase between 0.8 to 5.2 ppm in the complexes, between +9.8 to +15.3 in the transition structures, and further increase to between +16.2 and +54.9 ppm in the molecules. Spin−Spin Coupling Constants for Complexes, Transition Structures, and Molecules. In molecules and transition structures, the C4−C1 coupling constant is 1xJ(C4−C1) across

structures and NHC−CO2 molecules, the IR band corresponding to the symmetric stretch of CO2 is coupled to NHC bond stretches, making analyses complicated. However, this is not the case for the asymmetric stretch. The asymmetric stretching frequencies are reported in Table 6 and are not very sensitive to complex formation. Thus, this frequency is 2401 cm−1 in isolated CO2, and it lies between 2386 and 2403 cm−1 in the complexes. In the transition structures, the stretching frequency decreases to between 2114 and 2249 cm−1, and it further decreases to between 1760 and 1792 cm−1 in the NHC−CO2 molecules. Figure 9 shows the excellent correlation between the

Figure 9. C−O asymmetric stretching frequencies for complexes, transition structures, and molecules vs the C4−C1 distance.

values of the asymmetric stretching frequency and the C4−C1 distance for complexes, transition structures, and molecules, with a correlation coefficient of 0.998. Two distinctive O−C−O bending bands are observed in the IR spectra of complexes, one an in-plane bend and the other an out-of-plane bend. It is the in-plane bend that is more sensitive to complex formation. Its value of 659 cm−1 in the CO2

Table 7. Δδ13C1 and Δδ13C4 in Complexes, Transition Structures, and Molecules Δδ13C1,a ppm

a

Δδ13C4, ppm

NHC

NHC:CO2

NHC/CO2

NHC−CO2

NHC:CO2

NHC/CO2

NHC−CO2

δ13C4 of isolated NHCs

1 2 3 4 5 6 7

−0.73 −0.68 −0.63 −0.94 −0.70 −0.81 −0.82

−4.83 −6.41

−23.25 −20.66

15.34 12.65

54.87 42.80

−3.30 −3.51 −5.43 −4.48

−24.58 −24.46 −21.85 −25.15

4.32 2.44 1.11 5.14 5.23 3.43 0.80

9.84 11.60 9.97 −1.50

41.46 48.05 34.83 16.23

−8.43 −0.64 13.92 8.10 −5.84 9.18 34.08

The δ13C1 in isolated CO2 is 69.57 ppm. G


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Figure 11. 1t J(C4−C1), J(C4−O2), and J(C4−O3) vs the corresponding distances for NHC:CO2 complexes.

Figure 10. Change in the C1 chemical shielding upon complexation vs the C4−C1 distance.

order trendlines for J(C4−O2) and J(C4−O3) are 0.987 and 0.998, respectively, with J(C4−O2) for 7:CO2 removed. Why are the absolute values of J(C4−O3) greater than those of J(C4−O2) even though the C4−O3 distance is longer, and why was it necessary to remove the point for the complex 7:CO2 to obtain a good correlation between J(C4−O2) and the C4−O2 distance? As noted above, the CO2 molecule is tilted toward N5−H or C5−H in all of the complexes, and this tilt facilitates a secondary interaction that tends to localize the lone pair on O2 in this region. This is true for the ground state, and it should also be true in some of the excited states that couple to it through the FC operator. Thus, the lone pair of electrons on O2 is not as available for C4−O2 coupling as is the lone pair on O3. Moreover, the strength of this interaction is weakest in the complex 7-CO2, since it occurs between O2 and a C5−H group instead of N5−H. It appears from Figure 10 that 7-CO2 has too large a value of J(C4−O2) for its C4−O2 distance, since the electron pair on O2, which interacts with C5−H, is not as localized. The secondary interaction in complexes with NHCs 1−6 polarizes the O2 s electron density toward N5. This polarization reduces the s electron density at the O2 nucleus and reduces the absolute value of J(C4−O2). How do C4−C1 coupling constants change as complexes cross transition states to become molecules? 1tJ(C4−C1) for transition structures and 1J(C4−C1) for molecules are given in Tables 9 and 10, respectively. 1tJ(C4−C1) values for transition structures are negative, with values between −20 and −31 Hz. 1t J(C4−C1) values for these structures exhibit a second-order correlation with the C4−C1 distance, but the correlation coefficient is only 0.907. In contrast, 1J(C4−C1) values for molecules are positive, range from 49 to 57 Hz, and are plotted against the C4−C1 distance in the scattergram given as Figure S3 of the Supporting Information. The lack of a correlation is a

the tetrel bond; in molecules it is 1J(C4−C1) for the covalent C4−C1 bond. C4 is not bonded to O2 or O3 in complexes or transition structures, so the coupling constants are J(C4−O2) and J(C4−O3). In molecules, these carbon−oxygen coupling constants become two-bond couplings 2J(C4−O2) and 2J(C4− O3). The components of spin−spin coupling constants J(C4− C1), J(C4−O2), and J(C4−O3) in complexes, transition structures, and molecules are reported in Table S6 of the Supporting Information. The FC terms are excellent approximations to J(C4−C1) in all cases, J(C4−O2) and J(C4−O3) in complexes, and J(C4−O2) in transition structures. The FC terms do not approximate J(C4−O3) in transition structures as well, although the largest difference does not exceed 6%. The FC terms are not good approximations to 2J(C4−O2) and 2 J(C4−O3) in molecules, since the FC terms have reduced absolute values, while the PSO and SD terms have increased absolute values in molecules compared to transition structures. The values of 1tJ(C4−C1), J(C4−O2), and J(C4−O3) for complexes along with the corresponding distances are reported in Table 8. 1tJ(C4−C1) values are small and negative, lying between −1 and −4 Hz. Nevertheless, 1tJ(C4−C1) coupling constants exhibit a second-order dependence on the C4−C1 distance with a correlation coefficient of 0.993, as seen in Figure 11. What is also very interesting are the negative coupling constants J(C4−O2) and J(C4−O3) between atoms that are not directly bonded. For a fixed complex, the absolute value of J(C4−O3) is greater than J(C4−O2) except for 3:CO2, in which case they are equal, although the C4−O3 distance is greater than the C4−O2 distance in each complex. The distance dependence of these coupling constants can also be seen in Figure 11. The correlation coefficients of the second-

Table 8. 1tJ(C4−C1), J(C4−O2), and J(C4−O3) Coupling Constants (Hz) and Corresponding Distances (R, Å) for NHC:CO2 Complexes NHC

R(C4−C1)

1 2 3 4 5 6 7

2.876 2.939 3.001 2.792 2.848 2.897 2.838

1t

J(C4−C1)

R(C4−O2)

J(C4−O2)

R(C4−O3)

J(C4−O3)

−2.5 −1.7 −1.1 −4.1 −3.1 −2.4 −3.4

2.989 3.022 3.053 2.957 2.998 3.015 3.020

−5.8 −4.5 −3.4 −7.3 −5.7 −5.1 −8.2

3.299 3.362 3.427 3.220 3.266 3.311 3.227

−6.6 −4.8 −3.4 −9.5 −7.7 −6.3 −9.0

H


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Table 9. Spin−Spin Coupling Constants 1tJ(C4−C1), J(C4−O2), and J(C4−O3) (Hz) and Corresponding Distances (R, Å) for NHC/CO2 Transition Structures NHC

R(C4−C1)

1 2 4 5 6 7

2.113 1.990 2.285 2.246 2.083 2.225

1t

J(C4−C1)

R(C4−O2)

J(C4−O2)

R(C4−O3)

J(C4−O3)

−29.2 −31.0 −19.9 −21.2 −26.4 −20.8

2.566 2.498 2.672 2.659 2.562 2.673

−20.3 −20.3 −16.4 −15.4 −17.9 −21.9

2.720 2.637 2.844 2.816 2.704 2.772

−23.8 −23.5 −22.3 −21.6 −23.7 −23.5

Table 10. Spin−Spin Coupling Constants 1J(C4−C1), 2J(C4−O2), and 2J(C4−O3) (Hz) and Corresponding Distances (R, Å) for NHC−CO2 Molecules NHC

R(C4−C1)

1 2 4 5 6 7

1.524 1.535 1.534 1.539 1.543 1.542

1

J(C4−C1)

R(C4−O2)

55.1 48.9 56.5 52.3 50.5 53.7

2.303 2.299 2.295 2.298 2.296 2.329

2

J(C4−O2)

R(C4−O3)

−14.0 −14.9 −11.2 −10.6 −12.3 −13.7

2.303 3.009 2.347 2.353 2.341 2.329

2

J(C4−O3) −14.0 −15.0 −14.2 −15.2 −15.5 −13.7

reflection of the differences in the bonding in the NHC rings, which changes the electron distributions in the ground state and those excited states that couple to it through the Fermicontact operator. However, Figure 12 illustrates that C4−C1 coupling constants do show the evolution of the C4−C1 bond as complexes traverse transition structures to become molecules.

Figure 13. J(C4−O2) and J(C4−O3) for complexes and transition structures, and 2J(C4−O2) and 2J(C4−O3) for molecules vs the corresponding C4−O distances.

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Figure 12. 1tJ(C4−C1) for complexes and transition structures and 1 J(C4−C1) for molecules vs the C4−C1 distance.

CONCLUSIONS Ab initio MP2/aug′-cc-pVTZ calculations have been performed to identify equilibrium complexes and molecules and the transition structures that interconvert them, on the potential energy surfaces of a series of seven binary systems that have NHCs as the electron-pair donors to CO2. The results of these calculations support the following statements. 1. All of the NHCs form complexes with CO2 that are stabilized by C···C tetrel bonds, and six of the seven form stable molecules with C−C covalent bonds. 2. Six of the seven NHC:CO2 complexes are also stabilized by a secondary interaction between an N−H of the NHC and the adjacent O of CO2. The complex 7:CO2 has a C−H group adjacent to the O, and it is stabilized only by a tetrel bond. Because this secondary interaction is missing in 7:CO2, it has a smaller binding energy than 4:CO2, 5:CO2, and 6:CO2, even though 7 has the most negative MEP minimum.

J(C4−O2) and J(C4−O3) coupling constants for the transition structures are reported in Table 9, and 2J(C4−O2) and 2J(C4−O3) values for the molecules are reported in Table 10. These coupling constants remain negative and have their smallest absolute values in the complexes and their largest values in the transition structures. For the transition structures and the molecules with Cs symmetry, the absolute value of J(C4−O3) is greater than the absolute value of J(C4−O2). Neither J(C4−O2) and J(C4−O3) for transition structures, nor 2J(C4−O2) and 2J(C4−O3) for molecules, correlate with the corresponding C4−O2 and C4−O3 distances. The relationship between J(C4−O2) and J(C4−O3) and the changes in these coupling constants as a function of the C4− O distances as complexes cross transition states and become molecules can be seen in Figure 13. I



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ACKNOWLEDGMENTS This work was performed with financial support from the Ministerio de Economiá y Competitividad (Project No. CTQ2015-63997-C2-2-P) and Comunidad Autónoma de Madrid (S2013/MIT2841, Fotocarbon). Thanks are also given to the Ohio Supercomputer Center and CTI (CSIC) for their continued computational support.

3. Charge-transfer stabilizes all of the NHC:CO2 complexes and occurs from the NHC carbon atom to the antibonding in-plane π C−O bond of CO2. In the six complexes that have secondary stabilizing interactions, there is back-donation of charge from the O to the adjacent N−H group of the carbene. 4. Six of the seven NHCs form molecules with CO2 that are stabilized by C−C covalent bonds. With one exception, the molecules are significantly more stable than the corresponding complexes. 5. Transition structures present barriers to the interconversion of complexes and molecules. With one exception, the barrier for converting the molecule to the complex is much greater than the reverse barrier, since the molecule usually exists in a deeper potential well. 6. When the C−C distance in the transition structure is closer to the C−C distance in the complex, the barrier for converting the complex to the molecule is low. When the C−C distance in the transition structure is closer to that distance in the molecule, the barrier for converting the complex to the molecule is significantly higher. 7. AIM bonding parameters, shifts of IR C−O stretching and bending frequencies, and changes in NMR 13C chemical shieldings of the NHC carbon that is the electron-pair donor and the C of CO2 which is the electron-pair acceptor, are consistent with the changes that occur in the C···C tetrel bond in the complex as it goes through the transition state to become a covalent C−C bond in the molecule. 8. Spin−spin coupling constants 1tJ(C−C) and J(C−O) for complexes and transition structures, and 1J(C−C) and 2 J(C−O) for complexes, also dramatically illustrate the changing nature of the C−C bond as complexes go through transition structures to become molecules.

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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.7b08393. IRC reaction coordinates; structures, molecular graphs, and total energies for complexes, transition structures, and molecules; binding energies of N,N-dimethyl/CO2 complexes and molecules; plots of electron densities, Laplacians, and total energy densities at bond critical bonds versus the C−C distance; components of spin− spin coupling constants; scattergram for 2J(C−O) coupling constants versus the C−O distance for molecules (PDF)

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AUTHOR INFORMATION

Corresponding Authors

*E-mail: [email protected]. Phone: +1 330-609-5593. (J.E.D.B.) *E-mail: [email protected]. Phone: +34 915622900. (I.A.) ORCID

Janet E. Del Bene: 0000-0002-9037-2822 Ibon Alkorta: 0000-0001-6876-6211 José Elguero: 0000-0002-9213-6858 Notes

The authors declare no competing financial interest. J


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