Fostering the Basic Instinct of Boron in BoronâBeryllium Interactions

Article Cite This: J. Phys. Chem. A 2018, 122, 3313−3319

pubs.acs.org/JPCA

Fostering the Basic Instinct of Boron in Boron−Beryllium Interactions M. Merced Montero-Campillo,*,† Ibon Alkorta,† and José Elguero† †

Instituto de Química Médica (CSIC), Juan de la Cierva, 3, 28006 Madrid, Spain S Supporting Information *

ABSTRACT: A set of complexes L2HB···BeX2 (L = CNH, CO, CS, N2, NH3, NCCH3, PH3, PF3, PMe3, OH2; X = H, F) containing a boron− beryllium bond is described at the M06-2X/6-311+G(3df,2pd)//M0622X/6-31+G(d) level of theory. In this quite unusual bond, boron acts as a Lewis base and beryllium as a Lewis acid, reaching binding energies up to −283.3 kJ/mol ((H2O)2HB···BeF2). The stabilization of these complexes is possible thanks to the σ-donor role of the L ligands in the L2HB···BeX2 structures and the powerful acceptor nature of beryllium. According to the topology of the density, these B−Be interactions present positive laplacian values and negative energy densities, covering different degrees of electron sharing. ELF calculations allowed measuring the population in the interboundary B−Be region, which varies between 0.20 and 2.05 electrons upon switching from the weakest ((CS)2HB···BeH2) to the strongest complex ((H2O)2HB···BeF2). These B−Be interactions can be considered as beryllium bonds in most cases.

■

INTRODUCTION Ionic bonding and covalent bonding are usually described in terms of the electronegativity difference or similarity between the two partners involved in a common two center-two electron bond.1 Boron−beryllium bonds, formed by two electron−deficient atoms that usually behave as Lewis acids, were not described until 2015.2 In their seminal work, Arnold and collaborators took advantage of the nucleophilic character of the 1,3,2-diazaborolide anion for beryllium, proving the stability of such a link and describing the acceptor nature of beryllium in the resulting complex using electronic structure calculations. The ability of boron to behave as a Lewis base has recently attracted much attention, subverting its traditional Lewis acid role.3−12 Monovalent and trivalent boron donors are present in the cases reported in the literature, depending on what valence−shell structure is adopted by boron in the newly formed complex. The electronic structure of boron can in principle be modulated by the nature of the ligands surrounding it, and is evidenced by the geometrical features of the complex. Recently, Vondung and co-workers reported the case of borylenes (L2HB:) stabilized by phosphines,13 the latter acting as σ-donor and π-weakly acceptors, as examples of good boron donors toward metals. It is clear however that, for a neutral boron compound to act as an electron donor, a certain accumulation of charge is needed to foster its basic character, being equally important the good acceptor character of its counterpart in the Lewis pair. Beryllium is a powerful Lewis acid in closed-shell interactions,14,15 as revealed by the strength of beryllium bonds when compared to other closed-shell interactions.16,17 Just like hydrogen bonds, beryllium bonds are mainly of electrostatic © 2018 American Chemical Society

nature, but reinforced by a significant amount of charge transfer between a lone electron pair of the donor and the unoccupied σ*(BeX) and empty p(Be) orbitals of the BeX2 moiety acting as an acceptor. The strength of beryllium bonds results in high binding energies, an important distortion of the beryllium compound and a change in the intrinsic physicochemical properties of the donor due to the redistribution of the electron density.18−21 Therefore, if an appropriate nucleophilic boron compound is chosen, a beryllium bond can be formed between these two electron-deficient atoms. In this work, we explore the strength of this couple in a set of L2HB···BeX2 complexes, with L being small neutral molecules (L = HNC, CO, CS, N2, NH3, NCCH3, PH3, PF3, PMe3, H2O; X = H, F). As will be shown in the following sections, the nature of the boron valence shell is modulated by the ligands, but in all cases we found stable complexes in the gas phase due to the B → Be interaction.

■

COMPUTATIONAL DETAILS All geometries were optimized at the M06-2X/6-31+G(d) level of theory, a good compromise between computational cost and accuracy taking into account the size of some of the L ligands in the set. Harmonic frequency calculations allowed ensuring that the obtained structures were minima on the potential energy surface. The M06-2X functional combined with large basis sets has been shown to properly describe the energetic features of closed-shell interactions compared to CCSD(T) as a reference.22,23 For the estimation of the binding energies, we Received: February 13, 2018 Revised: March 14, 2018 Published: March 14, 2018 3313

DOI: 10.1021/acs.jpca.8b01551 J. Phys. Chem. A 2018, 122, 3313−3319

Article

The Journal of Physical Chemistry A

Figure 1. Optimized HB···BeX2 and L2HB···BeX2 complexes (L = CNH, CO, CS, N2, NH3, NCCH3, OH2, PH3, PF3, P(CH3)3; X = H, F) along with the corresponding B−Be distances (Å) and depart from planarity in the L2HB subunit (deg).Values in normal characters correspond to the BeF2 complexes (first row), whereas those corresponding to BeH2 complexes are represented in italics (second row).

The E(1)el term accounts for the electrostatic interaction energy of the unperturbed monomers, whereas the term E(1)exch corresponds to the first-order exchange energy contribution. The E(2)i term is the second-order induction energy resulting from the interaction between the permanent multipole and the induced multipole moments, and the consequent chargetransfer contributions plus the change in the repulsion energy due to the deformation of the monomers. The fourth term E(2)D, the second-order dispersion energy, is related to the multipole-induced multipole moment interactions plus the second-order correction for the coupling between the exchange repulsion and the dispersion interactions. Finally, a Hartree− Fock correction term (δHF) accounts for higher-order induction and exchange corrections, and then added to the induction energy. The DFT−SAPT calculations were carried out with the Molpro 2012 package.27 The nature of the B−Be interaction was analyzed through different computational approaches. The quantum theory of atoms in molecules (QTAIM)28−30 interprets bonding in a chemical system by looking at the topology of the electron density, allowing the localization of the nuclear attractors (NA) and the bonding paths between them. Along a given bond path, the density, laplacian of the density and energy density on the bond critical point (BCP) characterize the strength and nature of the bond. Also, other critical points of chemical interest such as ring critical points (RCPs) and cage critical points (CCPs) complete the topological description of a molecule. The QTAIM calculations were carried out with the AIMAll program.31 A different topological approach, the electron localization function (ELF),32 provides a Lewis-like picture of a molecule, allowing a partition of the space in basins associated with core, lone pairs (monosynaptic basins) and bonding

used single-point calculations at the M06-2X/6-311+G(3df,2pd)//M06-2X/6-31+G(d) level of theory. All these calculations were carried out with the Gaussian-09 program.24 The deviation from planarity of the HL2B subunits in their isolated or complexed forms is calculated by subtracting from 360° the values of all H−B−L and L−B−L angles. The binding energies are calculated as the energy difference between the optimized complex and the fully relaxed structures of the isolated compounds. This binding energy (BE) can be decomposed in two contributions, the interaction energy (Eint) and the deformation energy (Edef). The interaction energy is the difference between the energy of the complex and the sum of the energies of the interacting subunits within the geometry of the complex. The deformation energy for each subunit is the difference between the energy of this subunit in the geometry of the complex and the one of the fully relaxed structure. Therefore, for a given AB complex, the relationship between these magnitudes is the following: BE(AB) = E int(AB) + Edef (A) + Edef (B)

(1)

The density functional theory−symmetry adapted perturbation theory (DFT−SAPT)25,26 is used to investigate the nature of the interaction energy between two interacting compounds within the DFT framework. The energies of the monomers are then expressed in terms of orbital energies obtained from the Kohn−Sham density functional theory. This method allows estimating the contributions of the electrostatic, exchange, induction, and dispersion forces to the total interaction energy. The interaction energy is given by the sum of the following terms: E int = E(1)el + E(1)exch + E(2)i + E(2)D

(2) 3314


Article

The Journal of Physical Chemistry A regions (polysynaptic basins). The ELF function was calculated with TopMod.33 Finally, the Natural Bond Orbital (NBO) decomposition scheme34−36 provides an interpretation of bonding in terms of different types of localized orbitals (core orbitals, lone pairs, bonding orbitals, antibonding orbitals) and interactions between occupied and empty orbitals accounting for charge-transfer processes within the system. These calculations were carried out with the NBO 3.1 version included in the Gaussian 09 package.37

Table 1. Electronic Binding Energies, Interaction Energies, and Deformation Energies at the M06-2X/6311+G(3df,2pd)//M06-2X/6-31+G(d) level of theory (BE, Eint, Edef) for the HB···BeX2 and HL2B···BeX2 (L = CNH, CO, CS, N2, NH3, NCCH3, OH2, PH3, PF3, P(CH3)3; X = H, F) Complexes, Expressed in kJ/mola BE (E)

Eint

Edef

X H F

−104.3 −96.7

−149.9 −145.6

45.6 (0.9 + 44.7) 48.9 (0.8 + 48.0)

X H F H F H F H F H F F H F H F H F H F

−49.7 −68.4 −29.3 −41.8 −24.1 −37.7 −35.7 −47.2 −233.4 −263.0 −92.0 −128.3 −150.1 −59.6 −74.3 −162.1 −198.0 −252.5 −283.3

−82.1 −121.6 −48.4 −71.6 −38.6 −67.0 −58.9 −88.2 −320.3 −384.9 −207.6 −200.1 −244.1 −108.9 −135.7 −245.5 −311.3 −341.6 −403.9

32.3 (1.4 + 30.9) 53.3 (4.1 + 49.2) 19.0 (1.1 + 18.0) 29.7 (1.3 + 28.4) 14.5 (1.6 + 13.5) 29.3 (1.8 + 27.5) 23.3 (5.9 + 17.4) 41.0 (7.6 + 33.4) 86.9 (6.9 + 80.0) 121.9 (8.0 + 113.9) 115.6 (43.8 + 71.9) 71.9 (14.0 + 57.8) 94.0 (15.3 + 78.7) 49.3 (10.9 + 38.4) 61.4 (13.0 + 48.5) 83.4 (13.7 + 69.7) 113.3 (15.5 + 97.8) 89.1 (12.3 + 76.7) 120.5 (12.4 + 108.1)

HB···BeX2

■

RESULTS AND DISCUSSION In this section, we first describe how B−Be complexes look, paying attention to their main geometrical features. As a second step, the strength of the complexes is determined in terms of binding energies. We rationalize the different energy contributions to the binding energy looking at the interaction and deformation energy terms, and at the picture provided by the DFT−SAPT approach about the interaction energy itself. Finally, by means of the topology of the density and the NBO analysis, we characterize the nature of the B−Be interaction. Geometrical Features of Boron−Beryllium Complexes. The B−Be distance (Å) and the degree of deviation from planarity (deg) of the L2BH subunit in the equilibrium structures of the L2HB···BeX2 (L = CNH, CO, CS, N2, NH3, NCCH3, PH3, PF3, PMe3, OH2; X = H, F) complexes are represented in Figure 1. We provide the B−Be distance in reference monovalent boron complexes HB···BeH2 and HB··· BeF2. As reflected by the angle values in Figure 1, ligands such as CNH, CO, CS, and N2 do not essentially alter the planarity of the L2HB moiety (0.0°−7.3°), whereas complexes with donors such as NCCH3, PH3, and PF3 present a larger degree of pyramidalization (10.1°−14.3°), and the PMe3 complex is even more distorted (20.5°−21.4°). The largest departure from planarity is that of the four NH3 and OH2 complexes (28.0°− 52.5°). According to these geometries, from a formal point of view, we progressively change from a 2pz donor orbital to a hybridization for boron closer to sp3. It is important to note that the isolated (OH2)2HB compound is not planar (73.9°, see Table S1), meaning that only for this particular case, the complexation with BeX2 involves a decrease on the degree of pyramidalization. The shortest B−Be distances are precisely those from the ammonia and water complexes (1.891−1.923 Å), the largest distance being that of the L = CS, X = H case (2.386 Å). Shorter bonds correspond to larger departures from planarity (see Figure S1). Binding Energies and the Role of Deformation. We proceed to analyze the binding energies for the different {L, X} ligands to find out how strong boron−beryllium complexes are. Table 1 summarizes the binding energies (BE) in terms of electronic energy, providing also the contributions to the total BE arising from the interaction between subunits (Eint) and the deformation energy (Edef), as defined in the Computational Details section. In agreement with the geometries discussed in the previous section, the strongest complexes are (NH3)2HB··· BeX2 and (OH2)2HB···BeX2, with BEs between −233.4 and −283.3 kJ/mol, followed by the different phosphine complexes. The weakest complexes are those with CS as a donor ligand (−24.1/−37.7 kJ/mol). As derived from Table 1, all BeF2 complexes present larger BE values than the BeH2 ones. However, BeF2 is, in principle, a much better Lewis acid than BeH2, and consequently, the difference between interaction energies for a same couple is

HL2B···BeX2 L CNH CO CS N2 NH3 NCCH3 PH3 PF3 PMe3 OH2 a

Deformation energies include in parentheses the contributions from the boron and beryllium subunits, respectively.

much larger than in the case of the binding energies. The significant reduction between BeH2 and BeF2 complexes when looking at binding energies instead of interaction energies arises from much larger deformation energies in the BeF2 complexes. For instance, the (N 2 ) 2 HB···BeH 2 and (N 2 ) 2 HB···BeF 2 complexes present binding energies of −35.7 and −47.2 kJ/ mol, and interaction energies of −58.9 kJ/mol and −88.2 kJ/ mol. In this example, the deformation of the BeX2 moiety doubles upon going from X = H (17.4 kJ/mol) to X = F (33.4 kJ/mol). The very large interaction energies found for the water or ammonia complexes are also remarkably reduced by the huge amount of energy paid on deforming the interacting subunits from their fully relaxed geometries. Notably, in almost all cases, the main contribution to the deformation energy in the complexes is provided by the deformation of the beryllium moiety. Only (NCCH3)2HB···BeF2 presents a high contribution for the deformation energy arising from the boron subunit (43.8 kJ/mol), due to the quite different geometry presented by the methyl groups in the complex with respect to the free isolated (NCCH3)2HB compound. On the other hand, it is found that larger binding energies correspond to shorter B−Be distances. This is the expected trend for a wide range of binding energy values. Both distance and BE are exponentially correlated with R2 = 0.90, as shown in Figure 2. 3315


Article


Table 2. DFT Symmetry-Adapted Perturbation Theory (SAPT) Electrostatic (Eel), Exchange (Eexch), Induction (Ei), and Dispersion Terms (ED), in kJ mol−1a Eel HB···BeH2 HB···BeF2 HL2B···BeX2 L CNH

X H F

CO

H F

Figure 2. Exponential correlation between the B−Be distance (Å) and the binding energies (BE, kJ/mol) in absolute value for the L2HB··· BeX2 complexes (L = CNH, CO, CS, N2, NH3, NCCH3, OH2, PH3, PF3, P(CH3)3; X = H, F) at the M06-2X/6-311+G(3df,2pd)//M062X/6-31+G(d) level of theory.

CS

H F

N2

Forces Involved in the Interaction Energy. As explained above, the interaction energy is the main contribution to the binding energy, being in absolute value between 1.8 and 3.8 times larger than the deformation energy. Moreover, the binding and interaction energies are linearly correlated (R2 = 0.98, see Figure S2). In addition, we can estimate the amount of energy arising from the different forces involved in the interaction using SAPT (see Computational Details). Table 2 summarizes the values for the attractive and repulsive contributing terms within the SAPT approach, including in parentheses the percentage of the total attractive forces accounting for each attractive term (Eel, Ei, ED). The exchange energy is the largest term in absolute value for almost all complexes. Interestingly, the two strongest complexes in terms of binding energies, (NH3)2HB···BeF2 and (OH2)2HB···BeF2, have electrostatic terms slightly larger or almost equal to the exchange energy in absolute value. Within the attractive forces contributing to binding, in almost all cases, the electrostatic term is the dominant one, varying its contribution from a 37.1% (L = CS) to a 62.0% (L = OH2) of the total attractive terms. Note also that electrostatics is in general larger for BeF2 than for BeH2 complexes, whereas the dispersion contribution is always smaller in the BeF2 complexes. Complexes with linear ligands such as L = CNH, CO, CS and N2 present some particularities, in contrast to strongly interacting complexes. For instance, the dispersion terms are rather significant in this subset in comparison with the other complexes. In the (CS)2HB···BeH2 case in particular, the dispersion term represents a 25.6% of the attractive forces. Moreover, the induction energy in the (CS)2HB···BeF2 and (N2)2HB···BeF2 complexes is the dominant attractive term, beyond the electrostatic contribution. On the Nature of the Boron−Beryllium Interaction. The NBO decomposition scheme and the topology of the electron density provide a different picture of the nature of the B−Be interaction itself. The electron density at the B−Be BCP, the population of the ELF basin related to the B−Be interaction and the description of this bond provided by NBO are collected in Table 3. The B−Be distances are correlated with the topological parameters given in this table for the H2LB···BeX2 set (distance/ρ, exponential fit with R2 = 0.97, see Figure

H F

NH3

H F

NCCH3

F

PH3

H F

OH2

H F

Ei

ED

Eexch

−221.8 (50.1%) −192.0 (56.3%)

−160.8 (36.3%) −117.9 (34.6%)

−59.8 (13.5%) −31.2 (9.1%)

303.3

−127.0 (47.5%) −129.3 (49.6%) −83.7 (43.5%) −79.3 (44.2%) −54.7 (38.8%) −56.4 (37.1%) −85.7 (41.1%) −83.6 (40.8%) −445.1 (54.8%) −525.2 (61.3%) −190.6 (47.9%) −290.1 (54.7%) −307.4 (58.3%) −536.4 (55.3%) −628.4 (62.0%)

−91.3 (34.1%) −93.4 (35.8%) −68.9 (35.8%) −70.7 (39.4%) −50.0 (35.5%) −65.9 (43.3%) −81.9 (39.3%) −91.0 (44.4%) −256.4 (31.5%) −268.1 (31.3%) −161.8 (40.6%) −168.4 (31.7%) −169.6 (32.1%) −299.0 (30.8%) −311.9 (30.8%)

−49.2 (18.4%) −38.1 (14.6%) −39.7 (20.6%) −29.4 (16.4%) −36.1 (25.6%) −29.8 (19.6%) −40.9 (19.6%) −30.3 (14.8%) −111.4 (13.7%) −63.1 (7.4%) −45.9 (11.5%) −72.1 (13.6%) −50.7 (9.6%) −134.1 (13.8%) −74.0 (7.3%)

190.6

212.5

158.2 147.2 122.3 106.4 100.7 151.7 131.1 487.4 496.6 206.0 344.4 317.8 601.6 629.7

a The Hartree−Fock correction (δHF) is included in the Ei term, as explained in the Computational Details.

S3;38,39 distance/ELF population, linear fit with R2 = 0.96, see Figure S4). The corresponding molecular graphs are visualized in Figure S5 in the Supporting Information. As shown in Table 3, a BCP at the B···Be interatomic region is found in all cases. The electron density at this B−Be BCP increases from 0.020 (L = CS) to 0.080 au (L = OH2). The B− Be interaction is characterized by positive laplacian and negative energy density values (see Table S2 in the Supporting Information), the largest values corresponding to the shortest distances. Some representative examples of the ELF results are shown in Figure 3. Notably, linear molecules (CS, CO, CNH, N2) give place to small, monosynaptic V(B) polarized boron basins pointing toward the beryllium acceptor, with ELF populations from 0.20 to 1.03e. Electron-pair donors such as NH3, OH2, PH3 and PMe3 give place to disynaptic V(B,Be) bonds populated between 1.46 and 2.05e. The inductive effect triggered by the methyl groups on the P atom is accompanied by larger V(B,Be) ELF populations (1.58, 1.65e) than in the PH3 case (1.46, 1.54e), whereas the electron-withdrawing effect of the fluorine substituents in PF3 remarkably depopulates the interatomic B−Be region (1.07, 1.21e). According to the NBO description, the electrons involved in the B−Be interaction are mostly provided by the boron subunit 3316


Article


Table 3. Distances (Å), Electron Densities (ρ, e/bohr3) at the B−Be BCP, ELF Population of the Basin within the B−Be Interatomic Region, and B−Be NBO Second-Order Interaction Energies or Bonding Orbitals (BD) at the M06-2X/6311+G(3df,2pd)//M06-2X/6-31+G(d) level of theory d(B−Be)

ρ (e/ bohr3)

ELF (e)

HB···BeH2

1.876

0.066

2.11

HB···BeF2

2.002

0.051

2.00

2.269

0.029

0.47

HL2B···BeX2 L CNH

X H

F

CO

CS

NH3

0.030

0.59

H

2.273

0.025

0.54

F

2.240

0.027

0.64

H

2.386

0.020

0.20

F

N2

2.285

2.264

0.023

0.20

H

2.204

0.030

0.91

F

2.168

0.033

1.03

H

1.917

0.074

1.98

d(B−Be)

ρ (e/ bohr3)

ELF (e)

F

1.923

0.076

2.00

NCCH3

F

2.112

0.045

1.27

PH3

H

2.007

0.054

1.46

F

2.015

0.055

1.54

H

2.093

0.042

1.07

F

2.118

0.040

1.21

H

1.991

0.056

1.58

F

1.992

0.058

1.65

H

1.891

0.080

2.03

F

1.897

0.080

2.05

NBO BD (1.94e) 0.8393B (sp0.71) + 0.5436Be (sp2.92) BD (1.91e) 0.8780B (sp0.81) + 0.4786Be (sp2.36)

LP(B) > LP*Be 307.4 kJ/mol; BD(B−H) > LP*(Be) 34.7 kJ/mol LP(B) > LP*Be 284.5 kJ/mol; BD(B−H) > LP*(Be) 17.9 kJ/mol LP(B) > LP*Be 287.8 kJ/mol; BD(B−H) > LP*(Be) 53.8 kJ/mol LP(B) > LP*Be 290.5 kJ/mol; BD(B−H) > LP*(Be) 44.5 kJ/mol LP(B) > LP*Be 163.0 kJ/mol; BD(B−H) > LP*(Be) 46.1 kJ/mol LP(B) > LP*Be 204.0 kJ/mol; BD(B−H) > LP*(Be) 104.1 kJ/mol LP(B) > LP*Be 408.4 kJ/mol; BD(B−H) > LP*(Be) 64.7 kJ/mol LP(B) > LP*Be 448.4 kJ/mol; BD(B−H) > LP*(Be) 49.1 kJ/mol BD (1.93e) 0.8585B (sp1.52) + 0.5128Be (sp2.27)

PF3

PMe3

OH2

NBO BD (1.94e) 0.8738B (sp1.85) + 0.4864Be (sp1.37) LP(B) > LP*Be 718.0 kJ/mol; BD(B−H) > LP*(Be) 27.6 kJ/mol BD (1.71e) 0.9086B (sp4.04) + 0.4176Be (sp2.83) LP(B) > LP*Be 761.2 kJ/mol; BD(B−H) > LP*(Be) 14.5 kJ/mol LP(B) > LP*Be 640.5 kJ/mol; 2 BD(B−P) > LP* (Be) 37.9 kJ/mol; BD(B−H) > LP*(Be) 29.3 kJ/mol LP(B) > LP*Be 544.2 kJ/mol; 2 BD(B−P) > LP* (Be) 17.6 kJ/mol BD(B−H) > LP*(Be) 16.3 kJ/mol BD (1.76e) 0.9049B (sp3.24) + 0.4256Be (sp2.41) LP(B) > LP*Be 885.6 kJ/mol; 2 BD(B−P) > LP* (Be) 19.5 kJ/mol; BD(B−H) > LP*(Be) 17.5 kJ/mol BD (1.98e) 0.8445B (sp1.31) + 0.5356Be (sp2.30) BD (1.97e) 0.8618B (sp1.56) + 0.5072Be (sp1.47)

Figure 3. ELF (0.85) for HB:BeH2 and some selected L2BH···BeH2 complexes. Red lobes denote monosynaptic basins (lone pairs), green lobes denote disynaptic basins involving two heavy atoms, and yellow lobes are disynaptic basins involving hydrogen.

(see also the B and Be partial atomic charges in Table S3). Looking at the results in Table 3, most of the complexes are interpreted by this approach as a donation from a boron lone pair to an empty beryllium orbital, with very high interaction energies. Some lone pair donors acting as ligands in the boron subunit give place to a proper B−Be orbital, but with a strongly

polarized character and a much higher contribution from B with respect to Be, prevailing the donor role of boron. A second dative interaction identified by the NBO decomposition scheme is that of the bonding orbital associated with the B− H bond toward beryllium that, although smaller than the previous one, is also relevant in terms of energy. 3317



■

NBO and ELF descriptions coincide in giving similar pictures of the B−Be interaction, as the most populated V(B,Be) ELF basins appear together with proper bonding orbitals BD(B−Be) in the NBO approach, with populations near two electrons, as corresponds to typical polarized covalent bonds. Regarding the rest of the cases, taking into account the topological description provided by Table 3 and the slightly negative energy densities in Table S3, the closed-shell interactions between B and Be shown in this series present the standard features of beryllium bonds, i.e, significant binding energies, a certain amount of charge transfer and electron sharing from the donor to the empty Be orbital, and large deformation energies in the BeX2 subunit as a consequence of the interaction.

AUTHOR INFORMATION

Corresponding Author

*E-mail: [email protected]. ORCID

M. Merced Montero-Campillo: 0000-0002-9499-0900 Ibon Alkorta: 0000-0001-6876-6211 José Elguero: 0000-0002-9213-6858 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 (S2013/MIT2841, Fotocarbon). We also thank the ́ Centro de Computación Cienti fica de la Universidad Autónoma de Madrid (CCC-UAM) and CTI (CSIC) for their continued computational support.

■

CONCLUSIONS A series of L2HB···BeX2 complexes (L = CNH, CO, CS, N2, NH3, NCCH3, PH3, PF3, PMe3, OH2; X = H, F) has been studied with different computational techniques to analyze the strength and nature of B−Be interaction, with B and Be acting as Lewis base and Lewis acid, respectively. The geometry of the donor gradually changes from sp2 + pz to ∼ sp3. Water, ammonia and phosphine derivatives give place to the strongest complexes, with a maximum binding energy of −283.3 kJ/mol (L= OH2, X = F), whereas the weakest complexes are those with L = CS (BE = −24.1/−37.7 kJ/mol). In almost all cases, the beryllium moiety is the one contributing the most to the total deformation energy of the complex, affecting the total binding energy. According to SAPT, the exchange energy is the largest term in absolute value for almost all complexes, but the two strongest complexes in terms of binding energies, (NH3)2HB···BeF2 and (H2O)2HB···BeF2, have electrostatic terms slightly larger or almost equal to the exchange energy in absolute value. Within the attractive forces, the Coulombic term is the dominant one in almost all cases, but induction and dispersion are very relevant in particular for the complexes with linear donor molecules. A BCP is found for all complexes between B and Be, and ELF basins in the B−Be interatomic region are populated within a range of 0.20 to 2.05e. According to the NBO approach, the B−Be interaction can be described in most complexes as the sum of two dative contributions from a lone pair of B and the B−H bonding orbital toward Be. Therefore, in this work we extend the usual map of beryllium bonds with typical oxygen or nitrogen-containing Lewis bases to less common boron-containing Lewis bases acting as donors. All in all, small neutral molecules interacting with B foster its basic character toward Be, giving place to stable complexes formed between two electron-deficient atoms of the periodic table. These results should encourage the use of nucleophilic boron compounds to obtain new B−Be chemicals. Synthetic routes involving basic boron are starting to be relevant.40

■

Article

■

REFERENCES

(1) Pauling, L. The nature of the chemical bond. IV. The energy of single bonds and the relative electronegativity of atoms. J. Am. Chem. Soc. 1932, 54, 3570−3582. (2) Ewing, W. C.; Kramer, T.; Mies, J.; Schuster, J. K.; Arnold, T.; Braunschweig, H. Beryllium bis(diazaborolyl): old neighbors finally shake hands. Chem. Commun. 2015, 51, 737−740. (3) Bickelhaupt, F. M.; Radius, U.; Ehlers, A. W.; Hoffmann, R.; Jan Baerends, E. Might BF and BNR2 be alternatives to CO? A theoretical quest for new ligands in organometallic chemistry. New J. Chem. 1998, 22, 1−3. (4) Radius, U.; Bickelhaupt, F. M.; Ehlers, A. W.; Goldberg, N.; Hoffmann, R. Is CO a Special Ligand in Organometallic Chemistry? Theoretical Investigation of AB, Fe(CO)4AB, and Fe(AB)5 (AB = N2, CO, BF, SiO). Inorg. Chem. 1998, 37, 1080−1090. (5) Kusevska, E.; Montero-Campillo, M. M.; Mó, O.; Yáñez, M. One electron bonds in frustrated Lewis TPB ligands: Boron behaving as a Lewis base. Angew. Chem., Int. Ed. 2017, 56, 6788−6792. (6) Cid, J.; Gulyás, H.; Carbó, J. J.; Fernández, E. Trivalent boron nucleophile as a new tool in organic synthesis: reactivity and asymmetric induction. Chem. Soc. Rev. 2012, 41, 3558−3570. (7) Rozas, I.; Alkorta, I.; Elguero, J. Monohydride and Monofluoride Derivatives of B, Al, N and P. Theoretical Study of Their Ability as Hydrogen Bond Acceptors. J. Phys. Chem. A 1999, 103 (44), 8861− 8869. (8) Alkorta, I.; Soteras, I.; Elguero, J.; Del Bene, J. E. The boron− boron single bond in diborane(4) as a non−classical electron donor for hydrogen bonding. Phys. Chem. Chem. Phys. 2011, 13 (31), 14026− 14032. (9) Chou, S.-L.; Lo, J.-I.; Peng, Y.-C.; Lin, M.-Y.; Lu, H.-C.; Cheng, B.-M.; Ogilvie, J. F. Identification of Diborane(4) with Bridging B− H−B Bonds. Chem. Scien. 2015, 6, 6872−6877. (10) Alkorta, I.; Elguero, J.; Del Bene, J. E. Boron as an Electron− Pair Donor for B···Cl Halogen Bonds. ChemPhysChem 2016, 17, 3112−3119. (11) Del Bene, J. E.; Alkorta, I.; Elguero, J. B4H4 and B4(CH3)4 as Unique Electron Donors in Hydrogen−Bonded and Halogen− Bonded Complexes. J. Phys. Chem. A 2016, 120, 5745−5751. (12) Alkorta, I.; Elguero, J.; Del Bene, J. E. Borylene as an electron− pair donor for P···B pnicogen bonds. Struct. Chem. 2017, 28, 1419− 1427. (13) Vondung, L.; Frank, N.; Fritz, M.; Alig, L.; Langer, R. Phosphine−stabilized borylenes and boryl anions as ligands? Redox reactivity in Boron−based pincer complexes. Angew. Chem., Int. Ed. 2016, 55, 14450−14454. (14) Yáñez, M.; Sanz, P.; Mó, O.; Alkorta, I.; Elguero, J. Beryllium bonds, do they exist? J. Chem. Theory Comput. 2009, 5, 2763−2771.

ASSOCIATED CONTENT

S Supporting Information *

The Supporting Information is available free of charge on the ACS Publications website at DOI: 10.1021/acs.jpca.8b01551. Different geometrical parameters, Cartesian coordinates, correlations between several magnitudes, molecular graphs, topological results, and NBO atomic charges. (PDF) 3318


Article

The Journal of Physical Chemistry A (15) Albrecht, L.; Boyd, R. J.; Mó, O.; Yáñez, M. Cooperativity between hydrogen bonds and beryllium bonds in (H2O)nBeX2 (n = 1−3, X = H, F) complexes. A new perspective. Phys. Chem. Chem. Phys. 2012, 14, 14540−14547. (16) Mó, O.; Yáñez, M.; Alkorta, I.; Elguero, J. Modulating the strength of hydrogen bonds through beryllium bonds. J. Chem. Theory Comput. 2012, 8, 2293−2300. (17) Martín-Sómer, A.; Montero-Campillo, M. M.; Mó, O.; Yáñez, M.; Alkorta, I.; Elguero, J. Some interesting features of non−covalent interactions. Croat. Chem. Acta 2014, 4, 291−306. (18) Martín-Sómer, A.; Lamsabhi, A−M.; Mó, O.; Yáñez, M. The importance of deformation on the strength of beryllium bonds. Comput. Theor. Chem. 2012, 998, 74−79. (19) Montero-Campillo, M. M.; Yáñez, M.; Lamsabhi, A−M.; Mó, O. Spontaneous H2 Loss through the interaction of squaric acid derivatives and BeH2. Chem. - Eur. J. 2014, 20, 5309−5316. (20) Martín-Fernández, C.; Montero-Campillo, M. M.; Alkorta, I.; Yáñez, M.; Mó, O.; Elguero, J. Large proton affinity enhancements triggered by non−covalent interactions. Chem. - Eur. J. 2018, 24, 1971. (21) Montero-Campillo, M. M.; Sanz, P.; Mó, O.; Yáñez, M.; Alkorta, I.; Elguero, J. Alkaline−earth (Be, Mg, Ca) bonds at the origin of huge acidity enhancements. Phys. Chem. Chem. Phys. 2018, 20, 2413−2420. (22) Tama, R.; Mó, O.; Yáñez, M.; Montero-Campillo, M. M. Characterizing magnesium bonds: main features of a non-covalent interaction. Theor. Chem. Acc. 2017, 136, 36. (23) Burns, L. A.; Vázquez-Mayagoitia, A.; Sumpter, B. G.; Sherrill, C. D. Density-functional approaches to noncovalent interactions: A comparison of dispersion corrections (DFT-D), Exchange-hole dipole momento (XDM) theory, and specialized functionals. J. Chem. Phys. 2011, 134, 084107. (24) 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. Gaussian 09, revision D.01; Gaussian, Inc.: Wallingford, CT, 2013. (25) Jeziorski, B.; Moszynski, R.; Szalewicz, K. Perturbation Theory Approach to Intermolecular Potential Energy Surfaces of van Der Waals Complexes. Chem. Rev. 1994, 94, 1887−1930. (26) Misquitta, A. J.; Podeszwa, R.; Jeziorski, B.; Szalewicz, K. Intermolecular Potentials Based on Symmetry−Adapted Perturbation Theory with Dispersion Energies from Time−Dependent Density− Functional Calculations. J. Chem. Phys. 2005, 123, 214103. (27) Werner, H.-J.; Knowles, P. J.; Knizia, G.; Manby, F. R; Schütz, M.; Celani, P.; Györffy, W.; Kats, D.; Korona, T.; Lindh, R. et al. Molpro, 2012. (28) Becke, A. D.; Edgecombe, K. E. A simple measure of electron localization in atomic and molecular systems. J. Chem. Phys. 1990, 92, 5397−5403. (29) Bader, R. F. W. Atoms in Molecules: A quantum theory; Clarendon: Oxford, 1990. (30) Matta, C F.; Boyd, R. J., Eds. The quantum theory of atoms in molecules; Wiley-VCH: Weinheim, Germany, 2007. (31) Keith, T. A. AIMAll, version 13.01.27; T. K. Gristmill Software: Overland Park, KS, 2015. (32) Savin, A.; Nesper, R.; Wengert, S.; Fässler, T. F. ELF: The Electron Localization Function. Angew. Chem., Int. Ed. Engl. 1997, 36, 1808−1832. (33) Noury, S.; Krokidis, X.; Fuster, F.; Silvi, B. TopMod Package; Université Pierre et Marie Curie: Paris, 1997. (34) Foster, J. P.; Weinhold, F. Natural hybrid orbitals. J. Am. Chem. Soc. 1980, 102, 7211−7218. (35) Reed, A. E.; Weinhold, F. Natural bond orbital analysis of nearHartree−Fock water dimer. J. Chem. Phys. 1983, 78, 4066−4073. (36) Reed, A. E.; Weinstock, R. B.; Weinhold, F. Natural population analysis. J. Chem. Phys. 1985, 83, 735−746. (37) Glendening, E. D.; Reed, A. E.; Carpenter, J. E.; Weinhold, F. NBO, version 3.1.; Theoretical Chemistry Institute, University of Wisconsin: Madison, WI, 1998. (38) 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. (39) 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. (40) Pécharman, A.-F.; Colebatch, A. L.; Hill, M. S.; McMullin, C. L.; Mahon, M. F.; Weetman, C. Nat. Commun. 2017, 8, 15022.

3319


Fostering the Basic Instinct of Boron in BoronâBeryllium Interactions

Recommend Documents

Fostering the Basic Instinct of Boron in BoronâBeryllium Interactions