C Tetrel Bonds and CâC Covalent Bonds - ACS Publications

This is an open access article published under an ACS AuthorChoice License, which permits copying and redistribution of the article or any adaptations for non-commercial purposes.

Article pubs.acs.org/JPCA

Carbenes as Electron-Pair Donors To CO2 for C···C Tetrel Bonds and C−C Covalent Bonds Janet E. Del Bene,*,† Ibon Alkorta,*,‡ and José Elguero‡ †

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

‡

S Supporting Information *

ABSTRACT: Ab initio MP2/aug′-cc-pVTZ calculations have been carried out to identify stable complexes and molecules and the transition structures that interconvert them on the potential surfaces of ten singlet carbene bases acting as electron-pair donors to CO2. The carbene bases include cyclic C(NHCH)2 or NHC, C(NH2)2, an oxygen heterocyclic carbene C(OCH)2 or OHC, C(OH)2, C(CH3)2, cyclic C 3 H 2 , CCCH 2 , CCl 2 , CCH 2 , and CF 2 . Carbene:CO 2 complexes stabilized by C···C tetrel bonds have been found on all potential surfaces, whereas carbene−CO2 molecules stabilized by C−C covalent bonds have been found on eight surfaces. Three of these molecules have open structures with C2v symmetry, whereas the remaining have cyclic three membered C−O−C rings with Cs symmetry. The transition structures which connect the complex and the molecule are bound on three of the potential surfaces. Whether the transition structure is bound or unbound relative to the carbene and CO2 depends on the relationship among C−C distances at the three stationary points on the surface. Charge-transfer interactions stabilize carbene:CO2 complexes. The primary charge transfer in complexes arises from electron donation from the carbene lone-pair to the CO2 molecule. There is also back-donation of charge from CO2 to the carbene in three complexes. Systematic changes in bonding properties occur as complexes go through transition structures and become molecules. EOM-CCSD inter- and intramolecular C−C and C−O spin−spin coupling constants have been computed and compared for complexes and molecules. A search of the CSD database found the (NH2)2C−CO2 structure and 17 NHC− CO2 derivatives. Computed bond distances and angles have been compared with experimental data.

■

continue to be at the forefront of carbene research.10,11 The reactant that interacts with the carbenes in Scheme 1 is CO2, which can be viewed as a Jekyll and Hyde molecule. It is a major greenhouse gas, but it is also an attractive C1 building block12 that is abundant, cheap, nontoxic, and nonflammable.13 CO2 can be removed from the atmosphere by reaction with NHCs.14−16 There have been a few theoretical studies of reactions involving carbenes and CO2. Among these are two investigations of steric and solvent effects on the reaction of NHC derivatives with CO2.17,18 The very unstable carbene CH2 has been investigated in a study of the possible pathways for the reaction CH2 + CO2 → CH2O + CO, which indicated that the preferred path on the singlet surface evolves through oxiran-2one.19 A study of complexes of CH2 and CF2 with CO2 stabilized by C···C tetrel interactions has also been reported.20 In the present study, we have searched for stationary points on potential surfaces of ten different singlet carbenes acting as electron-pair donors to CO2. Three different types of stationary points have been found that correspond to a noncovalent complex stabilized by an intermolecular C···C tetrel bond, a

INTRODUCTION The reaction of carbenes with CO2 can produce compounds with two different structures depending on the nature of the carbene. Stable carbenes such as nitrogen heterocyclic carbenes (NHCs, 1) yield mesomeric betaines with a C−C bond between the carbene and CO2 (2).1 Very reactive carbenes (3) yield oxiranones (α-lactones) (4).2,3 These reactions are illustrated in Scheme 1. Among the carbenes, the N-heterocyclic carbenes (NHCs) have received the greatest attention. Several reviews4−6 and recent books7−9 have been published that discuss many of the important aspects of their reactivity. As a result, NHCs Scheme 1. Reaction of NHCs (1) and Reactive Carbenes R2C (3) with CO2 To Yield Betaines (2) and Oxiranones (4), Respectively

Received: April 10, 2017 Revised: May 3, 2017

© XXXX American Chemical Society

A

DOI: 10.1021/acs.jpca.7b03405 J. Phys. Chem. A XXXX, XXX, XXX−XXX

Article

The Journal of Physical Chemistry A

carbene:CO2 complexes using the NBO-6 program.36 Because MP2 orbitals are nonexistent, the charge-transfer interactions have been 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 some electron correlation effects. Carbon−carbon coupling constants across tetrel and covalent bonds, inter- and intramolecular carbon−oxygen coupling constants, and coupling constants across possible hydrogen bonds, in particular complexes and molecules, were evaluated using the equation-of-motion coupled cluster singles and doubles (EOM-CCSD) method in the CI (configuration interaction)-like approximation,37,38 with all electrons correlated. For these calculations, the Ahlrichs39 qzp basis set was placed on 13C, 15N, 17O, and 19F, and the qz2p basis set on 35Cl and those 1H atoms that might be involved in hydrogen bonding. The Dunning cc-pVDZ basis set was used for the remaining 1H atoms. The 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. The EOM-CCSD calculations were performed using ACES II40 on the HPC cluster Oakley at the Ohio Supercomputer Center.

molecule with a C−C covalent bond, and a transition structure that connects these two minima. The reactants and stationary points are depicted in Scheme 2. The carbenes included in this Scheme 2. Selected Stationary Points on the Potential Surface for the Reaction of a Singlet Carbene with CO2

study are a nitrogen heterocyclic carbene NHC, which is C(NHCH)2, C(NH2)2, an oxygen heterocyclic carbene C(OCH) 2 or OHC, C(OH) 2 , C(CH 3 ) 2 , cyclic C 3 H 2 , CCCH2, CCl2, CCH2, and CF2. In this paper we present and discuss the structures and binding energies of the complexes and molecules, and the transition structures that interconnect them on the potential surfaces. In addition, we also present and discuss the NBO properties of all stationary points, the chargetransfer energies of complexes, and spin−spin coupling constants for the equilibrium structures found on the surfaces.

■

RESULTS AND DISCUSSION Singlet carbenes act as electron-pair donors to CO2 to form complexes with C···C tetrel bonds and molecules with C−C covalent bonds. In the following sections we will first discuss the complexes, followed by the molecules, and then the transition structures that interconvert them. The next section presents a discussion of the bonding properties as complexes go through transition structures to form molecules. Spin−spin coupling constants for complexes and molecules are presented and discussed in the following section. The final section presents results of a search of the CSD database, and a comparison between computed and experimental bond distances and angles. Complexes. The structures, total energies, and molecular graphs of the ten stable carbene:CO2 complexes found on the potential surfaces are given in Table S1 of the Supporting Information. Figure 1 illustrates the complexes (OH)2C:CO2,

■

METHODS The structures of CO2 and the singlet carbene bases NHC, C(NH2)2, OHC, C(OH)2, C(CH3)2, cyclic C3H2, CCCH2, CCl2, CCH2, and CF2 were optimized previously21 at secondorder Møller−Plesset perturbation theory (MP2)22−25 with the aug′-cc-pVTZ basis set.26 This basis set was derived from the Dunning aug-cc-pVTZ basis set27,28 by removing diffuse functions from H atoms. Searches of the carbene:CO2 potential surfaces were then carried out at the same level of theory for complexes stabilized by C···C tetrel bonds, molecules with C−C covalent bonds, and transition structures that interconvert the complex and the molecule. Frequencies were computed to establish that the optimized structures correspond to equilibrium structures on their potential surfaces with no imaginary frequencies, and that transition structures have one imaginary frequency along the interconversion coordinate. Optimization and frequency calculations were performed using the Gaussian 09 program.29 The binding energies of the carbene:CO2 complexes, molecules, and transition structures were computed as the negative of the reaction energy (−ΔE) for the formation of these entities from the corresponding carbene and CO2. The electron densities of complexes, molecules, and transition structures have been analyzed using the Atoms in Molecules (AIM) methodology30−33 employing the AIMAll program.34 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 intermolecular bond critical point (ρBCP), the Laplacian (∇2ρBCP) at that point, and the total energy density (HBCP) have also been evaluated. The natural bond orbital (NBO) method35 has been used to obtain the stabilizing charge-transfer interactions for the

Figure 1. Complexes (OH)2C:CO2 with Cs symmetry, (CH3)2C:CO2 with C1 symmetry, and F2C:CO2 with C2v symmetry, illustrating the numbering system.

(CH3)2C:CO2, and F2C:CO2, which have Cs, C1, and C2v symmetry, respectively. In complexes with Cs or C1 symmetry, O2 is the O atom of CO2 that is closer to the carbene C4 atom. Figure 1 also indicates the numbering system that will be used in the following discussion. Structures and Binding Energies. The binding energies of the carbene:CO2 complexes and the C1−C4, O2−C4, C1−O2, and C1−O3 distances are reported in Table 1. The binding energies range from 8 to 25 kJ·mol−1 and decrease with respect to the carbene in the order B


Article


trendline, and that its binding energy is too large for its C1−C4 distance, a consequence of additional stabilization by the O5− H7···O2 hydrogen bond. The molecular graph of this complex indicates the presence of an O5−H7···O2 hydrogen bond as well as an interaction between O2 and C4. The molecular graphs of NHC:CO2 and (NH 2) 2 C:CO2 also indicate interactions between O2 and the N−H group of the carbene adjacent to it. In addition to these three complexes of Cs symmetry that are stabilized by secondary interactions, there is another complex, Cl2C:CO2, which also has this same symmetry. Cl2C:CO2 has a longer O2−C4 distance than the three other complexes with Cs symmetry, but a shorter O2−C4 distance than any of the complexes with C2v symmetry. The C1−C4 distances in the carbene:CO2 complexes range from 2.876 Å in the most stable complex NHC:CO2, to 3.163 Å in the least stable complex F2C:CO2. Both the C1−O2 and C1−O3 distances are essentially unchanged in the complexes with the smallest binding energies. The C1−O2 distance increases and the C1−O3 distance decreases but by no more than 0.005 Å in the three complexes with the largest binding energies. The complexes with C2v symmetry have two different equilibrium structures. In complexes with OHC and CCH2, the CO2 molecule lies in the symmetry plane defined by the carbene, whereas in those with cyclic C3H2, CCCH2, and CF2, the CO2 molecule is perpendicular to that plane, as illustrated in Figure 1 by F2C:CO2. However, the energy difference between the planar and perpendicular orientations of CO2 in optimized complexes is insignificant, because it is less than 2.0 kJ·mol−1. The equilibrium (CH3)2C:CO2 complex shown in Figure 1 has C1 symmetry. It is 0.56 kJ·mol−1 more stable than a complex with C2 symmetry which has 1 imaginary frequency. Charge-transfer energies are also reported in Table 1. The primary charge-transfer interaction in the most strongly bound carbene:CO2 complexes with Cs symmetry arises from electron donation from the C4 lone-pair to the antibonding C1−O2 orbital in the symmetry plane. These energies are 17.0, 19.4, and 7.4 kJ·mol−1 for complexes with NHC, C(NH2)2, and C(OH)2, respectively. There is also back-donation to the carbene from the O2 lone pair to the σ antibonding H−N5 or H−O5 orbital of the carbene, with charge-transfer energies of 2.3, 2.5, and 13.5 kJ·mol−1, respectively. These charge transfers are consistent with secondary electrostatic interactions in the complexes with NHC and C(NH2)2, and with the distorted O5−H···O2 hydrogen bond in the complex with C(OH)2. The complexes with C2v symmetry are also stabilized by charge transfer from the C4 lone pair of the carbene to the π antibonding O2−C1−O3 orbital of CO2 which lies in the symmetry plane of the carbene. The charge-transfer energies are 21.3 kJ·mol−1 for the complex with cyclic C3H2 and vary from 3.3 to 7.5 kJ·mol−1 for the remaining complexes. The complex with C(CH3)2 has C1 symmetry and a charge-transfer energy of 15.7 kJ·mol−1, with charge-transfer occurring from C4 to the σ antibonding O2−C1−O3 orbital. Molecules. There is a second minimum on all of the potential surfaces that can be described as a molecule with a C1−C4 covalent bond. However, the molecules formed from cyclic C3H2 and OHC are not stable relative to the isolated carbene plus CO2, so these have not been included in this section. There are two types of molecules, those that have an open structure, and those stabilized by a three membered C1− O2−C4 ring. Figure 3 illustrates the open structure of NHC− CO2 and the three membered ring structure of Cl2C−CO2.

NHC > C(NH 2)2 > C(OH)2 > C(CH3)2 > cyclic C3H 2 > OHC ≈ CCCH 2 > CCl 2 > CCH 2 > CF2

Table 1. Binding Energies (−ΔE) and Charge-Transfer Energies (ECT, kJ·mol−1), and C1−O2, O2−C4, C1−C4, and C1−O3 Distances (R, Å) in Carbene:CO2 Complexes carbene (sym)

−ΔE

R(C1− C4)

R(O2− C4)

R(C1− O2)a

R(C1− O3)a

ECTb

NHC (Cs) C(NH2)2 (Cs) C(OH)2 (Cs) C(CH3)2 (C1) cy C3H2 (C2v) OHC (C2v) CCCH2 (C2v) CCl2 (Cs) CCH2 (C2v) CF2 (C2v)

24.8 22.7 20.4 17.1 14.3 12.8 12.6 11.7 10.7 7.5

2.876 2.893 3.006 2.957 2.992 3.022 3.011 3.028 3.053 3.163

2.989 3.023 3.021 3.217 3.240 3.263 3.251 3.165 3.285 3.382

1.175 1.174 1.175 1.171 1.170 1.170 1.170 1.171 1.170 1.170

1.167 1.168 1.166 1.171 1.170 1.170 1.170 1.170 1.170 1.170

17.0c 19.4d 7.4e 15.7 21.3 7.5 7.2 6.3 5.8 3.3

a The C−O distance in the CO2 monomer is 1.170 Å. bClp → σ*O2− C1 orbital in the symmetry plane of complexes with Cs symmetry. Clp → π*O2−C1−O3 orbital in the plane of the carbene in complexes with C2v symmetry. Clp → σ*O2−C1−O3 orbital in the complex with C1 symmetry. cO2lp → σ*H−N5 orbital back-donation of 2.3 kJ/mol. d O2lp → σ*H−N5 orbital back-donation of 2.5 kJ/mol. eO2lp → σ*H−O5 orbital back-donation of 13.5 kJ/mol.

The three most stable complexes, NHC:CO2, C(NH2)2:CO2, and C(OH)2:CO2, have Cs symmetry and are stabilized by a C1···C4 tetrel bond and possibly by a secondary interaction between an N−H or O−H group of the carbene and O2, as can be seen in Figure 1 for the (OH)2C:CO2 complex. Whether these interactions are hydrogen bonds or simply electrostatic interactions can be determined by examining the H−N−O2 and H−O−O2 angles. The complexes of CO2 with NHC and C(NH2)2 have H−N−O2 angles of 42 and 40°, respectively, indicating that these are not N−H···O2 hydrogen bonds. In contrast, the H−O−O2 angle in (OH)2C:CO2 is 24°, compatible with a distorted, nonlinear O5−H7···O2 hydrogen bond. That this bond does contribute to the stability of the (OH)2C:CO2 complex can be seen in Figure 2, which is a plot of the complex binding energies versus the C1−C4 distance. It is evident that the point for (OH)2C:CO2 deviates from the

Figure 2. Binding energies of carbene:CO2 complexes (◆) and (OH)2C:CO2 (●) versus the C1−C4 distance. The exponential trendline has a correlation coefficient of 0.998. C


Article


molecules are reported in Table 2 and decrease with respect to the carbene in the order C(CH3)2 ≫ CCH 2 > linear CCCH 2 > CCl 2 ≫ CF2

The two most stable molecules relative to the carbene and CO2 have binding energies of 200 and 100 kJ·mol−1, respectively. These binding energies are significantly greater than the binding energies of the carbene−CO2 molecules with open structures. The high binding energy of the (CH3)2C− CO2 molecule reflects the relatively high instability of C(CH3)2. The molecules formed with linear CCCH2 and CCl2 have binding energies of about 75 kJ·mol−1. The least stable molecule is F2C−CO2 with a binding energy of 11 kJ·mol−1, which is only slightly greater than that of the complex at 8 kJ· mol−1. The atoms that form the three membered rings in these molecules are C1, O2, and C4. The C1−O2, O2−C4, and C1− C4 distances reported in Table 2 are distinctly different from the corresponding distances in the molecules with open structures. The C1−C4 distances vary between 1.41 and 1.45 Å and are therefore much shorter than these distances in molecules with open structures and covalent C1−C4 single bonds. The C1−O2 bond lengths vary from 1.34 to 1.40 Å and are significantly longer than the length of this bond in the molecules with open structures. This lengthening allows for the closing of the three membered ring. The O2−C4 distances in the complexes are between 1.44 and 1.50 Å, except for (CH3)2C−CO2. In this latter complex, it is 1.563 Å, which leads to distortion of the ring. The C1−O3 distances in the molecules are only slightly elongated relative to the distance in the CO2 monomer, because O3 is not involved in the ring. Because the variations in the C1−O2, O2−C4, and C1−C4 distances are coupled for the formation of the three membered ring, the binding energies of these complexes do not correlate with any one of these distances. Transition Structures and Energies. The transition structure that connects the complex and the molecule on the potential surface has been found on the eight surfaces. The structures, total energies, and molecular graphs of these are given in Table S3 of the Supporting Information. Their binding energies and C1−C4, C1−O2, and O2−C4 distances are given in Table 3. The binding energies have been computed consistent with the binding energies of the complexes and molecules. Thus, a binding energy that is positive refers to a transition structure that is stable relative to the isolated monomers, whereas a negative binding energy means that the

Figure 3. Carbene−CO2 molecules with (a) NHC, which has an open structure, and (b) CCl2, which has a cyclic C1−O2−C4 ring.

Table S2 of the Supporting Information reports the structures, total energies, and molecular graphs of the carbene−CO2 molecules. To differentiate between the molecule and the complex, we will write the molecule as carbene−CO2. Structures and Binding Energies of Molecules with Open Structures. Table 2 presents the binding energies of the three Table 2. Binding Energies (−ΔE, kJ·mol−1), and C1−O2, O2−C4, C1−C4, and C1−O3 Distances (R, Å) in Carbene− CO2 Moleculesa Molecules with Open Structures and C2v Symmetry −ΔE

carbene C(NH2)2 NHC C(OH)2

R(C1−C4)

R(C1−O2)b

R(O2−C4)

80.3 1.562 1.243 2.337 57.2 1.524 1.246 2.303 52.7 1.531 1.247 2.295 Molecules with C1−O2−C4 Rings and Cs Symmetry

carbene

−ΔE

R(C1− C4)

R(C1− O2)b

R(O2− C4)

R(C1− O3)b

C(CH3)2 CCH2 linear CCCH2 CCl2 CF2

207.5 105.4 78.4 72.3 10.8

1.449 1.406 1.412 1.451 1.444

1.340 1.396 1.378 1.357 1.383

1.563 1.449 1.492 1.502 1.444

1.195 1.185 1.186 1.186 1.184

a OHC and cyclic C3H2 do not form stable molecules with CO2. bThe C−O distance in the CO2 monomer is 1.170 Å.

molecules that have open structures of C2v symmetry. The most stable of these is (NH2)2C−CO2 with a binding energy of 80 kJ·mol−1, followed by NHC−CO2 and (OH)2C−CO2 with binding energies of 57 and 53 kJ·mol−1, respectively, relative to the carbene base and CO2. The H−N5−O2 angles that measure the nonlinearity of hydrogen bonds are 48 and 64° for (NH2)2C−CO2 and NHC−CO2, respectively, whereas the H− O5−O2 angle in (OH)2C−CO2 is 37°. These values indicate that there are no intramolecular hydrogen bonds in these molecules, although a stabilizing electrostatic interaction between O2 and N5−H9 or O5−H7 is possible. The molecular graphs of (NH2)2C−CO2 and (OH)2C−CO2 have bond paths consistent with such electrostatic interactions. The intramolecular C1−O2, O2−C4, C1−C4, and C1−O3 distances for these molecules are also reported in Table 2. The molecules with open structures have C1−C4 distances that are typical of single covalent C−C bond distances, with values between 1.52 and 1.56 Å. The C1−O2 distances are about 1.25 Å, an increase of 0.08 Å relative to the CO2 monomer. The O2−C4 distance is a long nonbonded distance of about 2.3 Å. Structures and Binding Energies of Molecules with Three Member Rings. There are five carbenes that form molecules with C1−O2−C4 rings. The binding energies of these

Table 3. Binding Energies (−ΔE, kJ·mol−1) and C1−C4, C1−O2, and O3−C4 Distances (R, Å) for Transition Structures carbenea

−ΔE

R(C1−C4)

R(C1−O2)

R(O2−C4)

NHC C(NH2)2 C(OH)2 C(CH3)2 linear CCCH2 CCl2 CCH2 CF2

9.3 13.7 −4.3 2.2 −28.3 −40.1 −26.1 −82.0

2.113 2.257 2.096 2.186 1.920 1.906 1.930 1.787

1.197 1.188 1.202 1.190 1.210 1.214 1.209 1.227

2.566 2.676 2.530 2.694 2.219 2.214 2.049 2.069

a

The carbenes are ordered according to decreasing binding energies of the complexes.

D


Article


Table 4. Values of δR1/δR2 (β) and the Binding Energies of Transition Structures (−ΔE, kJ·mol−1) for Carbene:CO2 Systems

transition structure is not bound relative to the monomers. The transition structures on the (OH)2C:CO2 and H2CCC:CO2 surfaces are illustrated in Figure 4.

a

carbene

βa

−ΔE

NHC (Cs) C(NH2)2 (Cs) C(OH)2 (Cs) C(CH3)2 (C1) linear CCCH2 (C2v) CCl2 (Cs) CCH2 (C2v) CF2 (C2v)

0.42 0.54 0.38 0.49 0.32 0.31 0.30 0.20

9.3 13.7 −4.3 2.2 −28.3 −40.1 −26.1 −82.0

See text for the definitions of δR1 and δR2.

is unbound, but by only −4.3 kJ·mol−1. For the remaining bases CCCH2, CCl2, CCH2, and CF2, β decreases to 0.32 or less, and the transition structures are unbound by −28 to −82 kJ·mol−1. In these latter systems, the transition state C1−C4 distance is much closer to the C1−C4 distance in the molecule, and this leads to a high transition state barrier. Bonding Properties of Complexes, Molecules, and Transition Structures. The AIM analyses provide three properties at each C1−C4 bond critical point (BCP), namely, the electron density (ρBCP), the Laplacian (∇2ρBCP), and the total energy density (HBCP). That these quantities reflect changes in the nature of the C1−C4 bond in going from the complex to the transition structure and then to the molecule can be seen in Figure S1 of the Supporting Information. The electron densities at bond critical points are near 0.01 au for the complexes, increase to between 0.05 and 0.13 au in the transition structures, and then further increase to between 0.20 and 0.30 au in the molecules. Once again, an excellent exponential relationship exists between the electron density at the bond critical point and the C1−C4 distance, with a correlation coefficient of 0.998.42−52 The values of the Laplacians initially increase with decreasing C1−C4 distance for the complexes, reach a maximum and then begin to decrease for transition structures, and continue to decrease and become negative for the molecules. The energy densities of the C1−C4 bonds are 0.001 au in the complexes, decrease to between −0.01 and −0.08 in the transition structures, and then further decrease to between −0.21 and −0.35 in the molecules. These changes are consistent with the changing nature of the C1−C4 bond, from a C1···C4 tetrel bond in the complexes to a shortened bond in the transition structures and then to a C1− C4 covalent bond in the molecules. Spin−Spin Coupling Constants for Complexes and Molecules. Table S4 of the Supporting Information provides the components of the carbon−carbon coupling constants 1t J(C1−C4) across tetrel bonds, the one-bond carbon−oxygen coupling constants 1J(C1−O2), and the intermolecular coupling constants J(O2−C4) for complexes. Table S5 gives the components of the one-bond coupling constants 1J(C1− O2) and 1J(C1−C4) for all molecules, 2J(O2−C4) for molecules with open structures, and 1J(O2−C4) for molecules with three membered rings. The coupling data for the (CH3)2C:CO2 complex are for the structure with C2 symmetry. Complexes. The coupling constants 1tJ(C1−C4), 1J(C1− O2), and J(O2−C4) are reported in Table 5. Values of 1tJ(C1− C4) are between 0.0 and −2.5 Hz over a range of 0.30 Å in the C1−C4 distance. This coupling constant tends to increase in absolute value as the C1−C4 distance decreases, but there is

Figure 4. Transition structures on the (OH)2C:CO2 and the H2CCC:CO2 surfaces.

Transition structures that are bound relative to the isolated molecules are found on the NHC:CO2, C(NH2)2:CO2, and (CH3)2C:CO2 surfaces, with binding energies of 9.3, 13.7, and 2.2 kJ·mol−1, respectively. The (OH)2C:CO2 structure is unbound, but by only −4.3 kJ·mol−1. The remaining transition structures are unbound relative to the corresponding monomers by −28 to −82 kJ·mol−1. From these energies, the computed barriers for the complex going to the molecule vary from 9 to 90 kJ·mol−1, whereas the reverse reactions going from the molecule to the complex have much higher barriers ranging from 48 to 205 kJ·mol−1. The only exception is the barrier on the F2C:CO2 surface, which is similar in both directions, because the binding energies of the complex and the molecule are similar. Why are some transition structures bound relative to the monomers and others unbound? Before answering this question, it should be noted that C1−C4, O2−C4, and C1− O2 distances in corresponding complexes, transition structures, and molecules are different, as can be seen by comparing these distances in Tables 1−3. In addition to distance changes, there are also changes in angular parameters. The CO2 molecule is linear in the complexes, but the O2−C1−O3 arrangement may be bent in transition structures and molecules. Another angular change is the orientation of the carbene molecule relative to CO2, as illustrated for the H2CCC:CO2 transition structure in Figure 4. However, although various bond lengths and angles change as complexes go through transition structures to become molecules, it is surely the C1−C4 distance which is the key structural parameter on the potential surface. This can be seen by computing δR1 = R(C1−C4)transition structure − R(C1−C4)molecule δR 2 = R(C1−C4)complex − R(C1−C4)molecule

β = δR1/δR 2 and then comparing β41 with the binding energies of complexes and molecules. The data required for this analysis are given in Table 4. If the transition structure were equidistant from the complex and molecule along the C1−C4 coordinate, then β would have a value of 0.5. Table 4 indicates that when β is between 0.42 and 0.54, the transition structure is bound relative to the corresponding isolated monomer and CO2. This occurs when the carbene bases are HNC, C(NH2)2, and C(CH3)2. β decreases to 0.38 for C(OH)2:CO2, and the transition structure E


Article


Table 5. C1−C4, C1−O2, and O2−C4 Coupling Constants (Hz) in Carbene:CO2 Complexes and Carbene−CO2 Molecules complexes carbene NHC C(NH2)2 C(OH)2 C(CH3)2 cyclic C3H2 OHC linear CCCH2 CCl2 CCH2 CF2

1t

J(C1−C4) −2.5 −2.5 −1.3 −1.7b −1.5 −1.2 −1.0 0.0 −0.4 −0.4

molecules

J(C1−O2)a

J(O2−C4)

19.0 19.0 20.0 19.4b 20.0 20.6 20.5 21.4 20.8 21.6

−5.8 −5.5 −3.7 −6.2b −5.6 −4.7 −4.1 −4.2 −3.2 −2.7

1

1

J(C1−O2)a

1

J(C1−C4)

2

J(O2−C4)

55.2 43.7 48.2 62.7

27.3 28.0 28.3 37.8

−14.0 −11.1 −11.4 2.7c

111.0 81.7 101.3 105.6

37.2 38.4 35.3 36.7

14.7c 14.9c 20.8c 19.4c

a1

J(C−O) for the CO2 monomer is 22.9 Hz at a C−O distance of 1.170 Å. bThe complex with C2 symmetry. c1J(O2−C4) values in molecules with ring structures

groups and the values of 1J(O2−C4) for these groups increase as the O2−C4 distance decreases, a good fit with a secondorder polynomial is almost guaranteed. What is most evident from Table 5 is the difference between corresponding coupling constants in complexes and molecules. 1 J(C1−C4) and 1J(C1−O2) in molecules are significantly greater than the absolute values of the corresponding coupling constants in complexes. With one exception, absolute values of 2 J(O2−C4) for molecules are greater than the absolute values of J(O2−C4) for complexes. CSD Search. A search of the CSD database found (NH2)2C−CO2 (Refcode: WETGAE)53 and 17 NHC−CO2 derivatives that are listed in Table S6. Reported in this table are experimental C1−C4, C1−O2, and C4−N bond distances and O2−C1−O3 and N−C4−N bond angles. The computed and experimental bond distances and angles for (NH2)2C−CO2, the computed bond distances and angles for NHC−CO2, and the average values of the corresponding distances and angles in the 17 nitrogen heterocyclic carbene molecules are reported in Table 6. The computed C1−O2 and C4−N bond distances are

scatter, as evident from the correlation coefficient of 0.828 for the second-order trendline relating 1tJ(C1−C4) to the C1−C4 distance. J(O2−C4) varies by only 3 Hz as the O2−C4 distance varies by 0.40 Å, but it is more negative than 1tJ(C1− C4). O2 and C4 are not bonded, and no correlation exists between J(O2−C4) and the O2−C4 distance. Complexation decreases 1J(C1−O2) from 22.9 Hz in the CO2 monomer to between 19.0 and 21.6 Hz even though the C1−O2 distance is unchanged in five complexes and increases by no more than 0.005 Å in the remaining complexes. Once again, there is no correlation between these two variables. The small values of J(N5−O2) and J(O5−O2) for the complexes with NHC, C(NH2)2, and C(OH)2 provide no evidence for the existence of hydrogen bonds in these complexes. Molecules. Table 5 reports 1J(C1−C4) and 1J(C1−O2) for all molecules, 2J(O2−C4) for those with open structures, and 1 J(O2−C4) for molecules with ring structures. 1J(C1−C4) values are between 44 and 55 Hz in the molecules with longer C1−C4 distances in open structures, and between 63 and 111 Hz in those with shorter distances in ring structures. The C1− C4 distances in the molecules (CH3)2C−CO2, Cl2C−CO2, and F2C−CO2 are within 0.007 Å, but 1J(C1−C4) values for these molecules are 63, 82, and 106 Hz, respectively. The remaining molecules with ring structures are H2CCC−CO2 and H2CC− CO2 with 1J(C1−C4) values of 111 and 101 Hz at C1−C4 distances of 1.412 and 1.406 Å, respectively. There is no correlation between 1J(C1−C4) and the C1−C4 distance in these molecules, and in general, such a correlation would not be expected because the covalent bonds formed by C4 in each molecule are different. Values of 1J(C1−O2) increase from 23 Hz in the CO2 molecule to about 28 Hz in molecules with open structures, and to between 35 and 38 Hz in those with ring structures. Although 1J(C1−O2) values for the five molecules with ring structures do correlate with the C1−O2 distance, the range of values for this coupling constant is quite small. The two-bond coupling constants 2J(O2−C4) are negative for the molecules with open structures. The one-bond coupling constants 1J(O2−C4) for molecules with three membered rings are positive and occur in three groups: H2CC−CO2 and F2C− CO2 have 1J(O2−C4) values of 20 Hz at an O2−C4 distance of 1.45 Å, H2CCC−CO2 and Cl2C−CO2 have 1J(O2−C4) equal to 15 Hz at a distance of 1.50 Å, and (CH3)2C−CO2 has 1 J(O2−C4) equal to 3 Hz at a distance of 1.56 Å. 1J(O2−C4) does correlate with the O2−C4 distance in these molecules, but because the coupling constants and distances occur into three

Table 6. Experimental and Computed Distances (R, Å) and Angles (∠, deg) of (NH2)2C−CO2 and NHC−CO2 Molecules (NH2)2C−CO2 exp R(C1−C4) R(C1−O2) R(C4−N) ∠O2−C1−O3 ∠N−C4−N

a

1.534 1.252 1.313 128.3 122.7

NHC−CO2 b

calc

exp

1.562 1.245 1.309 135.5 127.6

1.235 1.532 1.339 130.4 108.2

calc 1.246 1.523 1.335 135.8 106.2

a

The experimental values are for the CSD Refcode WETGAE. bThe experimental values are the average of 19 values in Table S6.

within 0.01 Å of the experimental values, whereas the computed C1−C4 bond distances are 0.03 and 0.01 Å longer than the values for (NH2)2C−CO2 and the average for the NHC−CO2 molecules, respectively. The computed O2−C1−O3 and N− C4−N angles in (NH2)2C−CO2 overestimate the experimental values by 7 and 5°, respectively. For NHC−CO2, the computed O2−C1−O2 angle overestimates the average of the experimental values by 5°, whereas the computed N−C4−N angle underestimates the average of the experimental values by 2°. The computed bond distances are in good agreement with the F


Article

The Journal of Physical Chemistry A experimental values, but the agreement between computed and experimental angles is not as good. However, when the quality of the agreement between theory and experiment is assessed, two factors should be kept in mind. First, the computed structures refer to isolated structures in a vacuum, whereas the experimental data are from crystal structures where packing considerations and intermolecular interactions play a role in structure determination. This can be seen in Figure 5 for

Figure 5. Hydrogen-bonded molecules of (NH2)2C−CO2 seen in Refcode WETGAE in the crystal.

WETGAE in which the (NH2)2C−CO2 molecules are arranged in hydrogen-bonded chains. Second, the NHC molecules found in the CSD have t-Bu or larger aromatic substituents bonded to the N atoms, whereas these bulky substituents have been replaced by hydrogen atoms in the computed NHC−CO2 molecule. No structures with an oxiranone group were found in the CSD database, most probably because these are highly reactive species that may rapidly be polymerized or decarboxylated.54

■

CONCLUSIONS Ab initio MP2/aug′-cc-pVTZ calculations have been carried out to identify stable complexes and molecules and the transition structures that interconvert them on the potential surfaces of ten carbene bases interacting with the CO2 molecule. The bases include the singlet carbenes cyclic C(NHCH)2 or NHC, C(NH2)2, cyclic C(OCH)2 or OHC, C(OH)2, C(CH3)2, cyclic C3H2, CCCH2, CCl2, CCH2, and CF2. The results of these calculations support the following statements. 1. Carbene:CO2 complexes exist on all potential surfaces and are stabilized by C···C tetrel bonds. In three of these, there are also secondary stabilizing interactions including a distorted hydrogen bond or electrostatic interactions. 2. Stable carbene−CO2 molecules have been found on all potential surfaces except those with cyclic C3H2 and OHC. Three of the stable molecules have open structures with C2v symmetry, whereas the remaining molecules are stabilized by cyclic three membered C1− O2−C4 rings. 3. The binding energies of the molecules are significantly greater than the binding energies of the complexes, except for F2C:CO2, for which the complex and molecule have similar binding energies. 4. The intermolecular C1−C4 distance is much longer in the complexes than in the molecules. The C1−O2 distance in complexes is short, but it elongates significantly in the molecules, particularly those stabilized by three membered rings. 5. Transition structures are bound relative to the isolated carbene and CO2 on the NHC:CO2, (NH2)2:CO2, and C(CH3):CO2 potential surfaces, slightly unbound on the C(OH)2:CO2 surface and are significantly unbound on the remaining surfaces. For these systems, the C1−C4 distance in the transition structure is much closer to the

■

distance in the carbene−CO2 molecule, and the resulting barrier to interconversion of the molecule and complex is high. Whether the transition structure is bound or unbound relative to the carbene and CO2 depends on the relationship involving the C1−C4 distances at the three stationary points on the surface. 6. Charge-transfer interactions stabilize carbene:CO2 complexes. The primary charge transfer arises from electron donation from the C4 lone-pair to CO2. There is also back-donation in three of these complexes from the O2 lone pair to the σ antibonding H−N5 or H−O5 orbital of the carbene. These back-donations are consistent with secondary hydrogen bonding and electrostatic interactions in these complexes. 7. Spin−spin coupling constants 1tJ(C1−C4) in complexes and 1J(C1−C4) in molecules do not correlate with the corresponding distances. 1J(C1−O2) and 1J(O2−C4) in molecules with ring structures do correlate with the corresponding distances. 1J(C1−C4) and 1J(C1−O2) in molecules are significantly greater than the values of the corresponding coupling constants in complexes. 8. A search of the CSD database found the (NH2)2C−CO2 structure and 17 NHC−CO2 derivatives. The computed bond distances are in good agreement with the experimental values, but the agreement between computed and experimental angles is not as good. However, given the different phases of the computed and experimental systems, and the interactions which occur in the solids, the agreement is quite acceptable. No structures with an oxiranone group were found in the CSD database.

ASSOCIATED CONTENT

S Supporting Information *

The Supporting Information is available free of charge on the ACS Publications website at DOI: 10.1021/acs.jpca.7b03405. Structures, total energies, and molecular graphs of complexes, transition structures, and molecules; plots of electron densities, Laplacians, and total energy densities at C(1)−C(4) bond critical points versus the C(1)− C(4) distance; components of spin−spin coupling constants; distances and angles in NHC−CO2 molecules from the CSD database (PDF)

■

AUTHOR INFORMATION

Corresponding Authors

*J.E.D.B.: e-mail, [email protected]; phone, +1 330-609-5593. *I.A.: e-mail, [email protected]; phone, +34 915622900, ORCID

Janet E. Del Bene: 0000-0002-9037-2822 Notes

The authors declare no competing financial interest.

■

ACKNOWLEDGMENTS This work was carried out with financial support from the Ministerio de Economiá y Competitividad (Project No. CTQ2015-63997-C2-2-P) and Comunidad Autónoma de Madrid (Project FOTOCARBON, ref S2013/MIT-2841). Thanks are also given to the Ohio Supercomputer Center and CTI (CSIC) for their continued computational support. G


Article


■

(24) Bartlett, R. J.; Silver, D. M. Many−Body Perturbation Theory Applied to Electron Pair Correlation Energies. I. Closed-Shell FirstRow Diatomic Hydrides. J. Chem. Phys. 1975, 62, 3258−3268. (25) Bartlett, R. J.; Purvis, G. D. Many-Body Perturbation Theory, Coupled-Pair Many-Electron Theory, and the Importance of Quadruple Excitations for the Correlation Problem. Int. J. Quantum Chem. 1978, 14, 561−581. (26) Del Bene, J. E. Proton Affinities of Ammonia, Water, and Hydrogen Fluoride and Their Anions: A Quest for the Basis-Set Limit Using the Dunning Augmented Correlation-Consistent Basis Sets. J. Phys. Chem. 1993, 97, 107−110. (27) Dunning, T. H. Gaussian Basis Sets for Use in Correlated Molecular Calculations. I. The Atoms Boron through Neon and Hydrogen. J. Chem. Phys. 1989, 90, 1007−1023. (28) Woon, D. E.; Dunning, T. H. Gaussian Basis Sets for Use in Correlated Molecular Calculations. V. Core-Valence Basis Sets for Boron Through Neon. J. Chem. Phys. 1995, 103, 4572−4585. (29) Frisch, M. J.; Trucks, G. W.; Schlegel, H. B.; Scuseria, G. E.; Robb, M. A.; Cheeseman, J. R.; Scalmani, G.; Barone, V.; Mennucci, B.; Petersson, G. A.; et al. Gaussian09, revision D.01; Gaussian, Inc.: Wallingford, CT, 2009. (30) Bader, R. F. W. A Quantum Theory of Molecular Structure and Its Applications. Chem. Rev. 1991, 91, 893−928. (31) Bader, R. F. W. Atoms in Molecules, A Quantum Theory; Oxford University Press: Oxford, U.K., 1990. (32) Popelier, P. L. A. Atoms In Molecules. An Introduction; Prentice Hall: Harlow, England, 2000. (33) Matta, C. F.; Boyd, R. J. The Quantum Theory of Atoms in Molecules: From Solid State to DNA and Drug Design; Wiley-VCH: Weinheim, Germany, 2007. (34) Keith, T. A. AIMAll, version 16.10.31; TK Gristmill Software: Overland Park, KS, 2011; aim.tkgristmill.com. (35) Reed, A. E.; Curtiss, L. A.; Weinhold, F. Intermolecular Interactions from a Natural Bond Orbital, Donor-Acceptor Viewpoint. Chem. Rev. 1988, 88, 899−926. (36) Glendening, E. D.; Badenhoop, J. K.; Reed, A. E.; Carpenter, J. E.; Bohmann, J. A.; Morales, C. M.; Landis, C. R.; Weinhold, F. NBO 6.0; Theoretical Chemistry Institute, University of Wisconsin: Madison, WI, 2013. (37) Perera, S. A.; Nooijen, M.; Bartlett, R. J. Electron Correlation Effects on the Theoretical Calculation of Nuclear Magnetic Resonance Spin−Spin Coupling Constants. J. Chem. Phys. 1996, 104, 3290−3305. (38) Perera, S. A.; Sekino, H.; Bartlett, R. J. Coupled−Cluster Calculations of Indirect Nuclear Coupling Constants: The Importance of Non−Fermi-Contact Contributions. J. Chem. Phys. 1994, 101, 2186−2196. (39) Schafer, A.; Horn, H.; Ahlrichs, R. Fully Optimized Contracted̈ Gaussian Basis Sets for Atoms Li to Kr. J. Chem. Phys. 1992, 97, 2571− 2577. (40) Stanton, J. F.; Gauss, J.; Watts, J. D.; Nooijen, M.; Oliphant, N.; Perera, S. A.; Szalay, P. S.; Lauderdale, W. J.; Gwaltney, S. R.; Beck, S.; et al. ACES II; University of Florida: Gainesville, FL, 1991. (41) Cioslowski, J. Quantifying the Hammond Postulate: Intramolecular Proton Transfer in Substituted Hydrogen Catecholate Anions. J. Am. Chem. Soc. 1991, 113, 6756−6760. (42) Knop, O.; Boyd, R. J.; Choi, S. C. Sulfur-Sulfur Bond Lengths, or Can a Bond Length Be Estimated from a Single Parameter? J. Am. Chem. Soc. 1988, 110, 7299−7301. (43) Gibbs, G. V.; Hill, F. C.; Boisen, M. B.; Downs, R. T. Power Law Relationships between Bond Length, Bond Strength and Electron Density Distributions. Phys. Chem. Miner. 1998, 25, 585−590. (44) Alkorta, I.; Barrios, L.; Rozas, I.; Elguero, J. Comparison of Models to Correlate Electron Density at the Bond Critical Point and Bond Distance. J. Mol. Struct.: THEOCHEM 2000, 496, 131−137. (45) Knop, O.; Rankin, K. N.; Boyd, R. J. Coming to Grips with N− H···N Bonds. 1. Distance Relationships and Electron Density at the Bond Critical Point. J. Phys. Chem. A 2001, 105, 6552−6566.

REFERENCES

(1) Kuhn, N.; Steimann, M.; Weyers, G. Synthesis and Properties of 1,3-Diisopropyl-4,5-Dimethylimidazolium-2-Carboxylate. A Stable Carbene Adduct of Carbon Dioxide [1]. Z. Naturforsch., B: J. Chem. Sci. 1999, 54, 427−433. (2) Wierlacher, S.; Sander, W.; Liu, M. T. H. Carboxylation of Carbenes in Low-Temperature Matrixes. J. Org. Chem. 1992, 57, 1051−1053. (3) Schröder, D.; Goldberg, N.; Zummack, W.; Schwarz, H.; Poutsma, J. C.; Squires, R. R. Generation of α-Acetolactone and the Acetoxyl Diradical •CH2COO• in the Gas Phase. Int. J. Mass Spectrom. Ion Processes 1997, 165/166, 71−82. (4) Herrmann, W. A.; Köcher, C. N-Heterocyclic Carbenes. Angew. Chem., Int. Ed. Engl. 1997, 36, 2162−2187. (5) Schuster, O.; Yang, L.; Rauberheimer, H. G.; Albrecht, M. Beyond Conventional N-Heterocyclic Carbenes: Abnormal, Remote, and Other Classes of NHC Ligands with Reduced Heteroatom Stabilization. Chem. Rev. 2009, 109, 3445−3478. (6) Hopkinson, M. N.; Richter, C.; Shedler, M.; Glorius, F. An Overview of N-Heterocyclic Carbenes. Nature 2014, 510, 485−496. (7) N-Heterocyclic Carbenes in Synthesis; Nolan, S. P., Ed.; WileyVCH: Weinheim, 2006. (8) N-Heterocyclic Carbenes: Effective Tools for Organometallic Synthesis; Nolan, S. P., Ed.; Wiley-VCH: Weinheim, 2014. (9) N-Heterocyclic Carbenes: From Laboratory Curiosities to Efficient Synthetic Tools, 2nd ed.; Diez-Gonzalez, S., Ed.; The Royal Society of Chemistry: Cambridge, U.K., 2017. (10) Hermann, M.; Frenking, G. Carbones as Ligands in Novel MainGroup Compounds E[C(NHC)2]2 (E = Be, B+, C2+, N3+, Mg, Al+, Si2+, P3+): A Theoretical Study. Chem. - Eur. J. 2017, 23, 3347−3356. (11) Georgiou, D. C.; Zhao, L.; Wilson, D. J. D.; Frenking, G.; Dutton, J- L. NHC-Stabilised AcetyleneHow Far Can the Analogy Be Pushed? Chem. - Eur. J. 2017, 23, 2926−2934. (12) Liu, Q.; Wu, L.; Jackstell, R.; Beller, M. Using Carbon Dioxide as a Building Block in Organic Synthesis. Nat. Commun. 2015, 6, 5933. (13) Solomon, S.; Plattner, G. K.; Knutti, R.; Friedlingstein, P. Irreversible Climate Change Due to Carbon Dioxide Emissions. Proc. Natl. Acad. Sci. U. S. A. 2009, 106, 1704−1709. (14) Vogt, M.; Bennett, J. E.; Huang, Y.; Wu, C.; Schneider, W. F.; Brennecke, J. F.; Ashfeld, B. L. Solid-State Covalent Capture of CO2 by Using N-Heterocyclic Carbenes. Chem. - Eur. J. 2013, 19, 11134− 11138. (15) Yang, L.; Wang, H. Advances in Carbon Dioxide Capture, Fixation, and Activation by Using N-Heterocyclic Carbenes. ChemSusChem 2014, 7, 962−998. (16) Maeda, C.; Miyazaki, Y.; Ema, T. Recent Progress in Catalytic Conversions of Carbon Dioxide. Catal. Sci. Technol. 2014, 4, 1482− 1497. (17) Denning, D. M.; Falvey, D. E. Substituent and Solvent Effects on the Stability of N-Heterocyclic Carbene Complexes with CO2. J. Org. Chem. 2017, 82, 1552−1557. (18) Denning, D. M.; Falvey, D. E. Solvent-Dependent Decarboxylation of 1,3-Dimethylimdazolium-2-Carboxylate. J. Org. Chem. 2014, 79, 4293−4299. (19) Kovacs, D.; Jackson, J. E. CH2 + CO2 → CH2O + CO, OneStep Oxygen Atom Abstraction or Addition/Fragmentation via αLactone? J. Phys. Chem. A 2001, 105, 7579−7587. (20) Alkorta, I.; Blanco, F.; Elguero, J.; Dobado, J. A.; Ferrer, S. M.; Vidal, I. Carbon···Carbon Weak Interactions. J. Phys. Chem. A 2009, 113, 8387−8393. (21) Del Bene, J. E.; Alkorta, I.; Elguero, J. Hydrogen-Bonded Complexes with Carbenes as Electron-Pair Donors. Chem. Phys. Lett. 2017, 675, 46−50. (22) Pople, J. A.; Binkley, J. S.; Seeger, R. Theoretical Models Incorporating Electron Correlation. Int. J. Quantum Chem. 1976, 10, 1−19. (23) Krishnan, R.; Pople, J. A. Approximate Fourth-OrderPerturbation Theory of the Electron Correlation Energy. Int. J. Quantum Chem. 1978, 14, 91−100. H


Article

The Journal of Physical Chemistry A (46) Alkorta, I.; Rozas, I.; Elguero, J. Molecular Complexes between Silicon Derivatives and Electron-Rich Groups. J. Phys. Chem. A 2001, 105, 743−749. (47) Knop, O.; Rankin, K. N.; Boyd, R. J. Coming to Grips with N− H···N Bonds. 2. Homocorrelations between Parameters Deriving from the Electron Density at the Bond Critical Point1. J. Phys. Chem. A 2003, 107, 272−284. (48) Espinosa, E.; Alkorta, I.; Elguero, J.; Molins, E. From Weak to Strong Interactions: A Comprehensive Analysis of the Topological and Energetic Properties of the Electron Density Distribution Involving X−H···F−Y Systems. J. Chem. Phys. 2002, 117, 5529−5542. (49) Alkorta, I.; Elguero, J. Fluorine−Fluorine Interactions: NMR and AIM Analysis. Struct. Chem. 2004, 15, 117−120. (50) Tang, T. H.; Deretey, E.; Knak Jensen, S. J.; Csizmadia, I. G. Hydrogen Bonds: Relation between Lengths and Electron Densities at Bond Critical Points. Eur. Phys. J. D 2006, 37, 217−222. (51) Alkorta, I.; Solimannejad, M.; Provasi, P.; Elguero, J. Theoretical Study of Complexes and Fluoride Cation Transfer between N2F+ and Electron Donors. J. Phys. Chem. A 2007, 111, 7154−7161. (52) Alkorta, I.; Elguero, J.; Del Bene, J. E. Pnicogen Bonded Complexes of PO2X (X = F, Cl) with Nitrogen Bases. J. Phys. Chem. A 2013, 117, 10497−10503. (53) Hervé, G.; Jacob, G. Novel Illustrations of the Specific Reactivity of 1,1-Diamino-2,2-Dinitroethene (DADNE) Leading to New Unexpected Compounds. Tetrahedron 2007, 63, 953−959. (54) L’abbé, G. Heterocyclic Analogues of Methylenecyclopropanes. Angew. Chem., Int. Ed. Engl. 1980, 19, 276−289.

I


C Tetrel Bonds and CâC Covalent Bonds - ACS Publications

Recommend Documents

C Tetrel Bonds and CâC Covalent Bonds - ACS Publications