Contrasting Behavior of the Z Bonds in XâZÂ·Â·Â·Y ... - ACS Publications

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Contrasting Behavior of the Z Bonds in X−Z···Y Weak Interactions: Z = Main Group Elements Versus the Transition Metals Jyothish Joy† and Eluvathingal D. Jemmis*,‡ †

School of Chemistry, Indian Institute of Science Education and Research-Thiruvananthapuram, Kerala, Thiruvananthapuram 695016, India ‡ Department of Inorganic and Physical Chemistry, Indian Institute of Science, Bangalore 560012, India S Supporting Information *

ABSTRACT: In contrast to the increasing family of weak intermolecular interactions in main-group compounds (X−Z··· Y, Z = main-group elements), an analysis of the Cambridge Structural Database indicates that electron-saturated (18electron) transition-metal complexes show reluctance toward weak M bond formation (X−M···Y, M = transition metal). In particular, weak M bonds involving electron-saturated (18electron) complexes of transition metals with partially filled dorbitals are not found. We propose that the nature of valence electron density distribution in transition-metal complexes is the primary reason for this reluctance. A survey of the interaction of selected electron-saturated transition-metal complexes with electron-rich molecules (Y) demonstrates the following: shielding the possible σ-hole on the metal center by the core electron density in 3d series, and enhanced electronegativity and relativistic effects in 4d and 5d series, hinders the formation of the M bond. A balance in all the destabilizing effects has been found in the 4d series due to its moderate polarizability and primogenic repulsion from inner core d-electrons. A changeover in the donor−acceptor nature of the metal center toward different types of incoming molecules is also unveiled here. The present study confirms the possibility of M bond as a new supramolecular force in designing the crystal structures of electron-saturated transition-metal complexes by invoking extreme ligand conditions. criteria for Z bond, though there exist exceptions,14 prompted Politzer et al. to unify all such weak interactions under the common roof of σ-hole interactions.15 The σ-hole interaction energy is found to be inversly proportional to the electronegativity of the Z atom.16 However, the recent discovery of weak interactions in noble gas compounds (aerogen bond)10a,17 and fluorine compounds18 led to the general anticipation that with suitable combination of X and Y groups any element in the periodic table (Z) can take part in weak bond formation. This general assumption appears to be valid in the main-group compounds. In contrast electronsaturated transition-metal complexes present a different story. Searches in the available primary literature and crystal structure database (CSD) gave many examples of the group 11 and 12 metals with weak interactions, especially the coinage metals with d10 configuration.19 Specifically, we could not find any standard 18-electron transition-metal complexes of the groups 3−10 where the metal is involved in a weak interaction of the X−M···Y kind. As a part of our continuing study of weak interactions we analyze the electronic structure of transitionmetal complexes with a view to find the factors that prevent

I. INTRODUCTION The last several decades witnessed flourishing noncovalent interactions across the periodic table similar to the ubiquitous hydrogen (H-bond, X−H···Y).1 Weak interactions (X−Z···Y) involving most of the elements (Z) of the main group have been identified.2−10 Following the H-bond,2 various terminologies are given to differentiate them such as lithium bond,3 beryllium bond,4 boron bond,5 tetral bond,6 pnicogen bond,7 chalcogen bond,8 halogen bond,9 aerogen bond,10 etc. The X− Z group in all of these examples is electron-saturated; specifically, the Z atom (group) has saturated its usual valence. A generic designation, the Z bond has been proposed for X− Z···Y interaction, where Z can be any element in the periodic table.11 The accepted requirement for weak interaction to be designated as Z bond is that both the X−Z and Y to be electron-saturated and that the Z should act as an electron acceptor during the X−Z···Y formation. This is obvious for the H-bond, with hydrogen’s one electron and one orbital. The energetics and directionality of the Z bonds are mostly decided by the σ-hole, a positive electrostatic potential region at the Z atom on the extension of X−Z bond.12 The stabilization here comes from the electrostatic attraction through the σ-hole along with reasonable magnitude of charge polarization at the donor−acceptor region.9b,13 The requirement of σ-hole as a © XXXX American Chemical Society

Received: August 28, 2016

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DOI: 10.1021/acs.inorgchem.6b02073 Inorg. Chem. XXXX, XXX, XXX−XXX

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Inorganic Chemistry

Figure 1. Deformation density maps indicating the contrast in the direction of charge polarization between Type I (M→Y, electron density depletion around the metal, red) and Type II (M←Y, accumulation of electron density around the metal, blue) complexes. Calculations are done at the M06-D3/QZ4P level of theory using ADF-2014. CSD reference code of the crystal structures (i) QUCBIA,30 (ii) EFUGOB,31 (iii) YOBRAJ,19b (iv) PEKSUT.32

ments.28 While these data support the reductionism we suggested earlier, where the definition of a chemical bond includes the H-bond, its more general version the Z bond, and the classical X−Y bond,16 we proceed here with the definition of M bond (X−M···Y), restricting X−M to be an electronsaturated transition-metal complex. Here, we seek reasons for the reluctance of electron-saturated transition-metal complexes to participate in noncovalent X−M···Y interactions compared to the electron-saturated main-group compounds. Model complexes with one and two metal atoms from group 9 are used in our analysis.

them from forming weak M bonds (X−M···Y interaction). Comparison to the main-group weak interactions helps in this. Our studies that suggest the possibilities for M bond are presented here. There are several examples in the literature where electronsaturated (18-electron) transition-metal complexes act as electron donor (Y) in X−Z···Y, where Z = hydrogen, halogen, etc.20 These also take the form of hemichelation and double hemichelation,21 donor−acceptor interactions,22 stabilizing heterodox bonding,23 and so on. Weak interactions in electron-unsaturated (16 or fewer electrons) transition-metal complexes present examples that are tantalizingly similar to the M···Y interactions and yet cannot be categorized as M bonds according to the definition. For example, the series (OC)5M··· Ng (M = Cr, Mo, W and Ng = Ar, Kr, Xe) showed remarkable structural stability with interaction energy ranging from 3 to 10 kcal/mol.24 The nature of continuum of bonding in these examples is obvious. If CO replaces Ng, we get a traditional M−L bond. It is possible to find a series of ligands that span a continuum of energy. A similar situation exists in the case of activation of small molecules, where σ-bond of the incoming molecule weakly interacts with the vacant coordination site of the electron-unsaturated transition-metal complex.25 Since this interaction involves the vacant coordination site of an electronunsaturated metal complex, the descriptive term weak dative bond (X−M←Y) is used in place of M bond (X−M···Y). Measurement of sequential bond dissociation energy for transition-metal cations (Mn+) with various ligands, Lx−1−M− L, also approaches the limit of noncovalent interaction when x is large.26 Intramolecular metal−ligand secondary bonding is another variety of noncovalent interactions observed in unsaturated transition-metal complexes.27 Machonkin and coworkers showed that such interactions (7 to 9 kcal/mol) are as strong as H-bonding via variable-temperature NMR measure-

II. RESULTS AND DISCUSSION A. Analysis of Cambridge Structural Database. In view of the large number of X-ray structures of transition-metal complexes available in the Cambridge crystallographic database, we reasoned that there might be examples where M bond is present. The ConQuest package available from CCDC is used to perform detailed searches in the Cambridge Structural Database (CSD; version 5.36/Conquest 1.17 November 2014). The initial search criterion was set to locate all possible transition metalany group (X) interaction at the van der Waal separation. We impose one restriction throughout the analysis that X should not be hydrogen. Preliminary searches resulted in a large number of structures with possible M bond in the crystal geometry. The search is refined further by specifying the desired geometry and coordination number. The structures are then sorted based on the total number of electrons in the complex to validate the electron saturation criteria (18-electron rule). Surprisingly, upon imposing the criteria of neutral complexes with 18-electrons, possible M bond candidates are found to be very few in the database. The majority of crystal structures obtained from the CSD are found to be unsaturated in the context of 18-electron rule. A few of B


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Inorganic Chemistry

Figure 2. Radial density distribution r2 R(r)2 of ns, np, nd, and (n + 1)s orbitals for (a) cobalt (n = 3), (b) rhodium (n = 4), and (c) iridium (n = 5) from scalar relativistic DFT calculations using ZORA-PBE/TZ2P level of theory. r = radial distance from the nucleus in angstroms.

electronic structure of transition-metal complexes in detail for a better understanding of the reluctance of Type I complexes to form weak M bonds. We begin with the concept of σ-hole, which helped to understand the nature of weak interactions in many well-known main-group noncovalent systems. According to Politzer et al.,12a,33 σ-hole on a halogen atom is generated when “a halffilled p-orbital on the halogen participates in forming a covalent bond, resulting in an electron deficiency in the outer (noninvolved) lobe of that p-orbital.” The magnitude of σhole on the Z atom of the X−Z molecule depends on the polarizability of the Z atom and electron-withdrawing power of the X group. Transition-metal complexes with lower oxidation states are soft acids and hence expected to show high polarizability, which in turn is anticipated to show strong σhole with proper combination of electron-withdrawing ligands. However, available literature data on Type I complexes does not support this. The apparent inability of transition-metal complexes to create σ-hole at the extension of X−M bond (X represents the collection of ligands to make the metal electron-saturated) involves geometrical constraints and electronic factors. Most common geometries in 18-electron complexes (octahedral, tetrahedral, and trigonal bipyramidal) possess the geometrical difficulty of having a vacant coordination site at the extension of X−M bond for the incoming Y group to attack. This can be solved by the proper design of electron-saturated complexes of square pyramidal geometry. The electronic factors originate from the d-orbitals coming into the valence space of the transition metals. We reason that the reluctance arises primarily due to the difference in the core and valence electron density distribution between main-group and transition-metal elements. The valence region of maingroup elements is composed of ns and np orbitals of comparable energy and spatial extension.34 In contrast, the ns and np orbitals of transition metals stay at the core level, and nd, (n + 1)s, and (n + 1)p orbitals span in the valence space. This changeover between the main-group and transition-metal elements has huge impact on the nature of electron density distribution at the van der Waals (vdW) surface of the molecule. Figure 2 shows the radial density distribution of ns, np, nd, and (n + 1)s orbitals of Co, Rh, and Ir. It is clear from Figure 2a that the valence 3d orbitals of Co place its radial density distribution inside the core 3s and 3p orbitals. It arises due to the lack of primogenic repulsion,35 that is, lack of repulsion from core orbitals of same angular distribution. The maximum

the selected electron-saturated crystal structures obtained from the CSD and corresponding Hirshfeld surface plots29 are shown in Figure S1. It clearly demonstrates the existence of weak interaction between the metal and the Y group. Most of the crystal structures retrieved from CSD this way belong to complexes of group 11 and 12 transition metals, Figure S2. This analysis led to the understanding that electron-saturated transition-metal complexes of groups 3−10 show an inherent reluctance to take part in weak M bond formation. It is in stark contrast to the nature of electron-saturated main-group compounds and group 11 and 12 transition-metal complexes. We selected nine complexes out of the structures obtained from CSD for further analysis. B. Electronic Structure Analysis of the Selected Complexes. The characteristic difference in the electronic structure of the nine crystal structures selected for further study is that five of them (Type I, Figure S3) are 18-electron complexes with partially filled metal d-orbitals, and four of them (Type II, Figures S3) are the ones with fully filled d-orbitals (d10). Figure 1 shows four of the selected crystal structures (two each from Type I and Type II) and its corresponding deformation density map (see Figure S4 for other complexes). The red and blue regions in the deformation density map correspond to electron density depletion and accumulation, respectively, during the complex formation from individual monomers. Hence, it essentially gives information about the direction of electron density flow (red→blue). Electron density difference (EDD) analysis on the selected crystal structures also gives the same information (Figure S5). To our surprise, the direction of charge polarization in Type I and Type II complexes exhibit opposite trend. In Type I complexes, the polarization direction is from the metal to the incoming molecule (Y), whereas in Type II it is from the incoming molecule to the electron-saturated metal. This analysis reveals a contrast in the direction of charge polarization among the complexes studied. Therefore, Type II complexes possess proper M bond characteristics, where the saturated metal center acts as an electron acceptor. However, Type I complexes failed to show the expected M bond behavior. Type II are electron-saturated complexes of group 11 and 12 transition metals. The diversity of weak M bonds in Type II complexes19 can be verified from the crystal structures given in Figure S2. These results are to be compared to the presence of weak Z bonds in almost all main-group elements including electronegative fluorine18 and especially to the absence of weak M bond in Type I complexes of electropositive transition metals. This realization has motivated us to analyze the C


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Inorganic Chemistry

Figure 3. MESP contour representation of the extent of σ-hole for main-group (F−Cl), Type I ((PF3)4RhI), and Type II ((PF3)3AgI) complexes. Blue and red contours represents positive and negative MESP, respectively. A cartoon representation is also given. Blue and red regions in the cartoon represent the σ-hole and core electron density, respectively. X represents the collection of ligands that make the system electron-saturated.

group (e.g., Co = 1.88, Rh = 2.28, and Ir = 2.20) as a consequence of relativistic effect.37 It also affects the polarizability of the metal center (Co = 7.5 Å3, Rh = 8.6 Å3, and Ir = 7.6 Å3).38 As a consequence of this heavier d-block elements also show reluctance toward weak M bond formation. A balance in all the effects can be seen in transition metals of the 4d series and hence the possible candidate for stronger M bonds. A comparison of molecular electrostatic potential (MESP) contours of Type I with Type II and main-group compounds (Figure 3) enables a general understanding of the nature of σhole across the periodic table. Relative position of the 0.04 atomic units (au) contour helps in this analysis (Figure S9). A cartoon representation of the net effect is also shown in Figure 3. The presence of core electrons at the valence region of 3d metals and increased electronegativity and relativistic effects in 4d and 5d metals results in diminished σ-hole strength in Type I complexes. However, significant radial separation between core and valence electrons in main-group and Type II compounds leads to stronger σ-holes. An equally important difficulty arises from the nature of higher occupied and lower unoccupied molecular orbitals (HOMOs and LUMOs) that are responsible for stabilization through charge polarization.39 The strong mixing of HOMOs and LUMOs under the influence of an effective polarizing agent (Y) is considered as charge-transfer interaction in the literature,40 which is one of the important stabilizing interactions in main-group Z bonds. In the case of 3d transition-metal complexes, metal contribution toward LUMOs is minimal, and hence stabilization due to the polarization would be small. As we go down the group, sd

probability region for the valence 3d electrons in cobalt are located at 0.35 Å away from the nucleus in comparison with 0.36 Å for 3s/3p core orbitals.34 It results in the shielding of the valence 3d orbitals by the core 3s/3p orbitals. As a consequence, the incoming ligands experience increased Pauli repulsion from the core electrons. This effect has been identified, and its implications are discussed in detail, for transition-metal complexes in many different contexts.36 For example, the elongated Mn−O bond in MnO4− ion is related to the higher Pauli repulsion experienced by the ligands from core 3s/3p orbitals that are located at the valence 3d region.36a This concept can be extended to describe the reluctance of Type I complexes to taking part in weak M bond formation in comparison with Type II complexes. In Type I complexes, the presence of core ns and np electrons effectively screen the valence nd orbital and hence block the possible existence of σhole at the metal center. It also makes the metal center more electron-rich and renders the vdW surface more prone to electrophilic attack. In Type II complexes, fully filled nd orbitals stay in the core region, and (n + 1)s and (n + 1)p orbitals are in the valence space. These (n + 1) s/p orbitals are highly diffused in space, and with proper combination of electron-withdrawing ligands σ-hole can be generated at the metal atom. Hence, Type II complexes behaves like a main-group Z-bond donor. As we go down the period, primogenic repulsion from the inner d-orbitals pushes the valence d-electrons out of the core region (Figure 2b,c for Rh and Ir).34 The radially diffused dorbitals in 4d and 5d series experience diminished screening from the inner core electrons and hence are anticipated to show stronger σ-holes. However, in contrast to main-group elements, electronegativity of the transition metals increases down the D


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Inorganic Chemistry

Figure 4. Model electron-saturated square pyramidal complexes of the Type I. 1 = Co(CN)2(NO)2Cl and 2M = M(CN)2(NO)2Mn(CO)5, M = Co, Rh, and Ir.

Figure 5. (a) Optimized geometry, (b) ELF plot, (c) electrostatic potential map at 0.001 au isosurface, and molecular orbitals ranging from HOMO−2 to LUMO+2 of the complex Co(CN)2(NO)2Cl. Calculations are done at the M06-D3/Def2-QZVP level of theory.

C. Design and Analysis of the M-Bond Complexes. Having analyzed the reasons behind the reluctance, we further endeavored to design a few model electron-saturated transitionmetal complexes to generate σ-hole at the metal center. The σholes are the positive potential region on the MESP map plotted over the vdW surface (0.001 au isosurface) of a molecule.41 It is logical to argue that the electron density distribution at the HOMOs determines the nature of electrostatic potential at the vdW surface. Hence, it is possible to generate σ-hole at the metal center by carefully engineering the frontier molecular orbitals (see “σ-Hole Generation by Orbital Engineering” section in the Supporting Information). Since we are going to use 18-electron square pyramidal

mixing increases, and metal contribution toward LUMOs enhances; however, energetic position of the LUMOs becomes unfavorable for effective overlap (Figure S10). In short, the hindrance toward electrostatic stabilization by the diminished σ-hole and orbital stabilization by the nature of the LUMOs cause the reluctance in Type I complexes to generate X−M···Y metal bonds. Though the presence of core electrons at the valence region is unavoidable, any strategy to generate M bonds should include ways to decrease the metal-based electron density at the HOMOs to strengthen the σ-hole and increase the coefficients over the LUMOs to enhance the chargepolarization effects. E


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Inorganic Chemistry

Table 1. Changes in the Selected Geometrical Parametersa of the Complexes Generated from 1 [Co(CN)2(NO)2Cl] and 2M [M(CN)2(NO)2Mn(CO)5] Optimized at the M06-D3/Def2-QZVP Level of Theory X−M···Y

M···Y (Å)

Δd (M−Lapical) (mÅ)

Δd (M−NO) (mÅ)

Δd (M−CN) (mÅ)

Δ∠(ON−M−NO)b (degree)

Δ∠(M−N−O)c (degree)

1←NCH 1···OCO 1···FCl 1···ClF 2Co···NCH 2Co···OCO 2Co···FCl 2Co···ClF 2Rh···NCH 2Rh···ClF 2Ir···NCH 2Ir···ClF

1.979 3.222 3.142 3.160 3.275 3.408 3.295 3.040 2.911 3.117 3.204 3.186

−5.11 −3.65 −1.03 0.95 −11.56 −6.93 −4.86 5.65 −15.74 5.50 −12.16 12.85

168.00 5.99 2.17 −1.62 3.14 1.10 0.65 −0.79 17.67 −2.29 5.29 −1.95

10.38 0.02 0.02 0.82 0.53 0.10 0.03 0.34 1.74 1.44 0.76 0.42

14.94 3.01 0.79 −2.28 0.24 0.21 0.08 −0.12 4.15 −0.09 0.68 −0.36

−32.848 −2.91 −0.88 +2.59 −0.89 −0.28 −0.97 +1.77 −5.74 +1.61 −2.32 +2.2

a M···Y = noncovalent distance, Δd (M−Lapical) = change in metal-apical ligand bond length, Δd (M−NO) and Δd (M−CN) = changes in metal− basal ligand bond length, Δ∠(N−M−N) and Δ∠(M−N−O) = change in ON−M−NO and M−N−O bond angles, respectively. M−Lapical = Co−Cl in complex 1 and M−Mn(CO)5 in complexes of 2M. bAngle is positive when ∠(N−M−N)dimer is greater than ∠(N−M−N)monomer. cAngle is positive when ∠(M−N−O)dimer is greater than ∠(M−N−O)monomer.

σ-hole (+34.95 kcal/mol) at the extension of the Co−Cl apical bond toward the vacant coordination site as indicated by the red arrow in the MESP map (Figure 5c). The π-donating apical Cl ligand perturbs the HOMO and HOMO−1 electron density away from the metal, and strong π-accepting NO+ and CN− ligands stabilize HOMO−2 with minimal lone-pair density at the metal center. Molecular orbial coefficients of reasonable magnitudes over the metal atom in the LUMOs (LUMO has 18% metal 3d contribution, and LUMO+2 has 29% 3d and 3% 4s metal contributions) could also encourage the stabilization through charge polarization. The electron localization function (ELF),48 a map of electron pair probability, plotted over the plane containing Co−Cl bond, Figure 5b, also indicates the existence of σ-hole at the metal center (see Figures S7 and S8 for details). Further, dipole moment of 2.9 D pointing toward the apical ligand from the metal center aids the formation of M bond with the negative end of an electron-rich Y group. The designed hetero-bimetallic complexes, 2M, possess σhole of reasonable magnitude at the vacant coordination site of the metal atom (Figures S11 and S12). Complex 2Co shows σhole strength of +12.4 kcal/mol at the metal center with dipole moment of 2.6 D. Influence of primogenic repulsion from 3d orbitals and higher polarizability of Rh (8.6 Å3) increase the σhole strength in complex 2Rh to +16.4 kcal/mol. However, it decreases to +13.9 kcal/mol in complex 2Ir due to the decreased polarizability (7.6 Å3) of Ir. Compared to complex 1, weaker σ-hole in complex 2Co is perhaps due to the lower electron-withdrawing ability of the apical ligand Mn(CO)5, even though we observed a marginal decrease in the molecular orbital coefficient over the HOMOs (Figure S11). Incoming molecules (Y) of varying donor strengths are used to probe the nature of weak interactions in complexes 1 and 2M. Two types of incoming molecules (Y) are used in the present study, (i) molecules with negative potential at the interacting atom (NCH, OCO, FCl) and (ii) molecule with positive potential over the interacting atom (ClF). It enables us to show the flexibility of the electron-saturated complexes to act both as an electron acceptor and donor upon weak bond formation. Complexes formed with electron-donating molecules belong to the M bond category, whereas electronaccepting ClF molecule forces the electron-saturated metal center to act as an electron donor leading to halogen bonding.

complexes for the following analysis, and its HOMO is a lone pair orbital projecting toward the vacant coordination site,42 any strategy to decrease this accumulated electron density would be desirable to generate σ-hole at the metal. We have arrived at a few model electron-saturated square pyramidal complexes as shown in Figure 4. Studies are performed using cobalt, rhodium, and iridium as the central metal atom, with strong π-acceptor ligands such as NO+ and CN− at the basal position, and chlorine/metal fragment at the apical position. All the model complexes are optimized at the M06-D3/Def2-QZVP level of theory using Gaussian 09 quantum chemistry package and are found to be minima on the potential energy surface.43 The complexes shown in Figure 4 prefer square pyramidal geometry as the global minimum over the traditionally most stable trigonal bipyramidal geometry of five-coordinated complexes. This geometrical preference is due to the stabilization provided by the strong π-acceptor ligands at the basal position. The optimized geometries show slight tilt and bend in the NO ligands from the square plane, and reasons for this observation have been discussed extensively.44 The Rh and Ir counterparts of the model complex M(CN)2(NO)2Cl (1) are not considered further because of the uncertainty in electron counting arising from the large deviation in Rh−N−O bond angle (140°) in Rh(CN)2(NO)2Cl complex45 and the preference for trigonal bipyramidal geometry as the global minimum for Ir(CN)2(NO)2Cl complex. Another step in this direction is to replace the apical Cl− ligand of Co(CN)2(NO)2Cl by suitable metal fragment to create electron-saturated hetero-bimetallic complexes of the type M(CN)2(NO)2−Mn(CO)5, M = Co (2Co), Rh (2Rh), and Ir (2Ir), as shown in Figure 4. Such complexes of middle and late transition metals are known to possess metallophilic interactions46 in extended one-dimensional structures. For example, dimers of hetero-bimetallic complex [PtM(tba)4(OH2)] (M = Co and Ni; tba = thiobenzoate) exhibit Pt···Pt metallophilic interactions, where anti-ferromagnetic exchange occurs through four dz2 orbitals.47 Such interactions are mostly stabilized by dispersion and relativistic effects. The optimized geometry, electrostatic potential map, and the nature of frontier molecular orbitals of the complex Co(CN)2(NO)2Cl (1) are shown in Figure 5. This complex shows F


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Inorganic Chemistry

Figure 6. QTAIM topological map and MESP map of the dimeric structures of the complex Co(CN)2(NO)2Cl (1) with various incoming molecules. (a) 1←NCH, (b) 1···OCO, (c) 1···FCl, (d) 1···ClF. Color coding in the MESP map ranges from red (maximum negative) to blue (maximum positive).

The noninnocent nature of NO ligand49 is used to differentiate weak noncovalent interactions from that of strong dative bonds. The linear arrangement of NO is expected to retain as in the parent complex for noncovalent M···Y interactions with NO+ configuration (NO acts as three electron donor45). However, it attains a bent geometry with NO− configuration (NO becomes one electron donor45) for strong M←Y dative bond, where bending in NO creates unsaturation at the metal center and encourages further electron acceptance leading to strong interaction. Changes in the relevant geometrical parameters of the resultant complexes as a consequence of the interaction with Y groups are given in Table 1. Very short M···Y distance and significant deviation in geometrical parameters upon 1←NCH formation reveal the nature of interaction as strong dative bonding. The minimal changes in the bonding parameters for all other compounds ensures the nature of interaction as weak noncovalent. The characteristic geometrical features shown in Table 1 can differentiate M bonding from halogen bonding in the complexes of 1 and 2M. Reduction in M−Lapical bond length (up to 15 mÅ) and enhancement in M−Lbasal bond length (up to 17 mÅ) are unique features of M bonding, and the reverse is observed in halogen bonding. Similarly, expansion and reduction in ON−M−NO and M−N−O bond angles, respectively, are characteristic of M bonding. Halogen bonding shows the opposite trend. The quantum theory of atoms in molecules (QTAIM) is employed to study the topological features of weak-bonding in the complexes studied here.50 QTAIM topological parameters of the M bonds are given in Table S1. The values of electron density (ρ) and Laplacian of density (▽2ρ) at the M···Y bond

critical point (BCP) are coming under the criteria proposed by Koch and Proplier for hydrogen bonding (ρ = 0.002 to 0.04 and ▽2ρ = 0.15 to 0.02)51 confirms the topological similarities between weak M bonding and hydrogen bonding. The positive value of ▽2ρ indicates the closed-shell nature of interaction at the M···Y region. The delocalization index δ(M···Y), which is a measure of the average number of electrons delocalized between the interacting atoms, is also in agreement with the characteristics of main-group Z bonding. 52 The BCP connecting central metal atom and incoming Y group visually demonstrates the existence of M bond, Figure 6. The MESP maps shown in Figure 6 reveal the electrostatic nature of the interaction, where blue region over the metal center and green/yellow area over the Y group denotes the σhole and negative potential regions, respectively. Complexes 1···OCO and 1···FCl in Figure 6b,c represents the proper alignment of the monomers for better electrostatic attraction, where negative end of the Y group points toward the σ-hole on the metal. However, a changeover in the nature of interaction can be seen in 1···ClF (Figure 6d), where the σ-hole over the chlorine atom of ClF molecule is pointing toward the σ-hole on the metal center. The existence of 1···ClF is surprising, because it is stable even in the presence of repulsive σ-hole−σ-hole interaction. It is also interesting to note that the σ-hole−σ-hole interaction in 1···ClF possesses comparable noncovalent distances with M bonded 1···FCl. QTAIM analysis and ESP maps of hetero-bimetallic complexes 2M···Y also shows similar behavior, Figure S13. The direction of electron density flow and visual identification of the nature of interaction in weak M bonds can be achieved by inspecting the deformation density map and noncovalent interaction (NCI) plot shown in Figures G


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Inorganic Chemistry Table 2. Energy Decomposition Analysis of the Studied Complexesa complex

ΔEPauli (kcal/mol)

ΔEElectrostatic (kcal/mol)

ΔESteric (kcal/mol)

ΔEOrbital (kcal/mol)

ΔEDispersion (kcal/mol)

binding energy (kcal/mol)

1←NCH 1···OCO 1···FCl 1···ClF 2Co···NCH 2Co···OCO 2Co···FCl 2Co···ClF 2Rh···NCH 2Rh···ClF 2Ir···NCH 2Ir···NCH (SR-ZORA) 2Ir···ClF 2Ir···ClF (SR-ZORA)

69.2 1.2 1.4 4.6 1.9 0.7 0.8 6.6 7.1 4.8 3.8 4.2 5.5 6.1

−58.1 (59.77%) −2.1 (52.50%) −1.6 (47.06%) −0.5 (6.85%) −2.7 (51.93%) −1.0 (47.48%) −0.8 (36.36%) −2.8 (26.41%) −7.3 (59.84%) −1.2 (15.00%) −4.5 (56.96%) −4.5 (55.56%) −1.5 (17.04%) −2.0 (20.62%)

11.1 −0.9 −0.2 4.1 −0.8 −0.3 0.0 3.8 −0.2 3.6 −0.7 −0.3 4.0 4.1

−36.7 (37.76%) −0.7 (17.50%) −0.8 (23.53%) −4.8 (65.75%) −1.1 (21.15%) −0.3 (13.04%) −0.4 (18.19%) −6.1 (57.55%) −3.4 (27.87%) −5.1 (63.75%) −1.9 (24.05%) −2.1 (25.92%) −5.6 (63.64%) −6.0 (61.86%)

−2.4 (2.67%) −1.2 (30.00%) −1.0 (29.41%) −2.0 (27.40%) −1.4 (26.92%) −1.0 (43.48%) −1.0 (45.45%) −1.7 (16.04%) −1.5 (12.29%) −1.7 (21.25%) −1.5 (18.99%) −1.5 (18.52%) −1.7 (19.32%) −1.7 (17.52)

−28.0 −2.8 −2.0 −2.7 −3.3 −1.6 −1.4 −4.0 −5.1 −3.2 −4.1 −3.9 −3.3 −3.6

a

At the M06-D3/QZ4P level of theory using ADF-2014 program. Percentage contributions to the total attractive interactions are given in parentheses. SR-ZORA is used to invoke the relativistic effects in Ir complexes. Steric term is the sum of Pauli and electrostatic components.

can be seen from Figure 7. It makes the metal center slightly more positive and enhances the interaction with the apical

S14 and S15, which unequivocally represents the characteristics of genuine M bond. The existence of the dimers 1···ClF and 2M···ClF is found to be counterintuitive in view of the occurrence of direct σhole−σ-hole interaction. The peculiar ability of the designed complexes to bind FCl molecule through both fluorine (electron-rich) and chlorine (electron-deficient) faces suggest the amphoteric behavior of the metal center; that is, the metal center can act both as Lewis acid and Lewis base depending on the nature of the incoming Y group. Energy decomposition analysis (EDA) is employed to gain detailed insight into the nature of M bonding and to the amphoteric behavior of the designed complexes, Table 2. We use the EDA method developed independently by Morokuma and by Ziegler and Rauk as implemented in the ADF package. Here the electrostatic component (ΔEelectrostatic) of interaction energy is treated quasi classically, and hence the effect of charge polarization is accounted for in the orbital term (ΔEorbital).53 Hence, our EDA description of the M bond mainly relies on the electrostatic and orbital components of the interaction energy as given in Table 2. EDA differentiates strong dative bonding in 1←NCH from that of the weak M bond in the rest of the complexes. The binding energy for dative bonding (−28.0 kcal/mol) is much larger than that for weak M bonding (−1.4 to −5.1 kcal/mol). The dative bond is stabilized by electrostatic and orbital components with negligible contribution from the dispersion, whereas the weak M bonds are stabilized mostly by electrostatic and dispersion components with minimal contribution from the orbital interaction. The diminished orbital component and enhanced dispersion interaction differentiates weak M bonds from maingroup Z bonds, where orbital interaction plays a major role in stabilizing the weak contact.54 The M−Lapical bond length shortening during the M bond formation is perhaps due to the enhanced dispersion and reduced orbital interaction. Analysis of the polarizability components along the three axes of the complexes 1 and 2M showed that it is higher along the x- and y-directions (i.e., the plane containing M−NO and M−CN bonds) compared to the z-axis (M−Lapical bond axis). Hence, during the M bond formation, the higher polarizability of M− NO bonds (e.g., 19.23 Å3 for Co−NO vs 12.65 Å3 for Co−Cl in complex 1) tend to disperse the electron density toward it, as

Figure 7. Changeover in the direction of charge polarization for M bonding (1···FCl) and halogen bonding (1···ClF). Red and blue region represent the depletion and accumulation of electron density, respectively. Calculations are done at the M06-D3/QZ4P level of theory.

ligand, which ultimately results in shortening of the M−Lapical bond. The resultant accumulation of extra electron density at the M−Lapical bond is evident from the increase in electron density (ρBCP) at the M−Lapical BCP (e.g., 15.74 mÅ bond length shortening in complex 2Rh···NCH goes with an increase of the ρBCP from 0.030 to 0.032 au). In contrast to the M bonding, halogen bond strength decreases as the σ-hole strength increases, where stronger σhole in 2Rh makes the 2Rh···ClF complex weaker than 2Co··· ClF and 2Ir···ClF. Adding scalar relativistic zeroth-order corrections (SR-ZORA) in 2Ir decreases the M bond strength by 0.2 kcal/mol in 2Ir···NCH and increases the halogen bond strength by 0.3 kcal/mol in 2Ir···ClF. It occurs primarily due to the decrease in σ-hole strength by the inclusion of relativistic effects. Hence, we conclude that the absence of primogenic repulsion in 3d series and presence of relativistic effects in 5d H


Article

Inorganic Chemistry series reduces the σ-hole strength and hinders the weak M bond formation. On the basis of the above analysis, we suggest that electron-saturated transition-metal complexes of the 4d series are better candidates for M bonding. It is evident from the EDA analysis that the ability of central metal atom to reverse the electrostatic and orbital components on demand is the primary reason for the occurrence of amphoteric behavior, Table 2. For example, the direct σhole−σ-hole interaction in 1···ClF leads to negligible electrostatic component (6.85%) with enhanced orbital interaction (65.75%) as compared to 1···FCl where 47.06% electrostatic and 23.53% orbital interaction exists. This effect is directly reflected in the steric component which is the sum of Pauli and electrostatic terms. Murray et al. showed that the counterintuitive stabilization of noncovalent interaction between σholes occurs mainly due to polarization/dispersion between the interacting fragments.55 Here we also notice that, along with stabilization due to dispersion, charge polarization (orbital component) also plays the crucial role. EDA-NOCV (NOCV = natural orbital for chemical valence) deformation density maps are helpful in demonstrating the changeover in the direction of electron density flow. The deformation density map of the complexes 1···FCl and 1···ClF are shown in Figure 7, where electron density flow direction is from red to blue. In the M bond complex 1···FCl, electron flow direction is from FCl molecule to the metal center, whereas the reverse occurs in halogen-bonded complex 1···ClF. Halogen bond formation involves the removal of electrons from M−Lapical bond, which results in bond length elongation as shown in Table 1. The above analysis led to the conclusion that even though there exists σ-hole at the metal center that takes part in weak M bond formation, the inherent tendency of these complexes is to act as electron donors rather than acceptors. Hence, the two-way interaction due to the amphoteric nature of the metal center diminishes the binding energy, and thus the absence of M bonding in crystal structures can be justified.

interaction plays a major role in stabilizing the complexes studied here. Our findings buttress the fact that dispersion alone cannot stabilize the M bond unless it is supported by electrostatic attraction from the σ-hole. The present study shows that a balance in all the effects (σ-hole screening, nature of LUMOs, polarizability of the metal, and the two-way bonding) is achieved in electron-saturated transition-metal complexes of the 4d series. These ideas are sure to attract the attention of experimentalists to come up with ligand designs to generate σ-hole and M bonds that act as a new class of supramolecular force in electron-saturated transition-metal complexes.

IV. COMPUTATIONAL DETAILS Geometry optimization of all complexes studied here are performed using Gaussian 09 quantum chemistry package at M06-D3/Def2QZVP level of theory.43 The M06 functional was shown to perform superior for transition-metal complexes and noncovalent systems, and hence it was used throughout this study.56 Very tight convergence criteria and ultrafine grid are used to locate the minimum-energy geometry. The use of quadruple-ζ quality basis set minimizes the errors due to basis set superposition.57 Vibrational frequencies are calculated to examine the true minimum nature of the optimized structures. Some complexes show one small negative frequency of the order less than 15 cm−1 due to the very shallow potential energy surface at the interaction region. EDA of the Morokuma−Ziegler type58 is performed to gain a detailed understanding of the mechanism of M bonding using ADF-2014 program package at M06-D3/QZ4P level of theory.59 EDA is capable of decomposing the binding energy at the M···Y region into various physically meaningful quantities like Pauli/exchange repulsion, electrostatic interaction, orbital interaction, and dispersion interaction.60 Orbital component of the total binding energy has further been decomposed into pairwise donor−acceptor interactions using ETS-NOCV scheme implemented in the ADF-2014 package.61 NOCV deformation density analysis gives detailed insight into the direction and magnitude of electron density reorganization during the M bond formation.62 QTAIM topological parameters are calculated using AIMALL package with the wave function files generated from the M06/Def2-QZVP calculations.63 MESP maps are visualized using the GaussView software, and EDD maps are computed using the Multiwf n program.64

III. CONCLUSIONS We have analyzed the reasons behind the reluctance of electron-saturated transition-metal complexes of partially filled metal d-orbitals (Type I) to take part in weak M bond formation. The partial screening of the possible σ-hole on the metal center by the core electrons and diminished charge polarization due to the minimal orbital coefficient on the LUMOs are the main reasons why M bonds of Type I complexes are rare. However, highly diffused valence (n + 1)s and (n + 1)p orbitals in Type II complexes (fully filled dorbitals) can sustain the σ-hole and form M bonds just like the main-group compounds. Hence we conclude that it is difficult to observe σ-hole at the metal center if d-orbitals in the valence space are partially occupied. A few electron-saturated transitionmetal complexes with extreme ligand conditions have been designed to generate σ-hole at the metal center to substantiate this. These complexes are allowed to interact with electron-rich donor molecules (Y) to form weak M bonds (X−M···Y). These M bonds are stabilized primarily due to electrostatic and dispersion interactions with minimal orbital component. We also found that the designed complexes could undergo counterintuitive σ-hole−σ-hole interaction with electron acceptor groups. A close inspection revealed that these can act both as Lewis acid and Lewis base by switching the relative magnitude of electrostatic and orbital components. Decomposition of the interaction energy showed that dispersion

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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.inorgchem.6b02073. QTAIM topological parameters, details of the CSD survey, designing σ-hole by orbital engineering and Cartesian coordinates of the studied complexes (PDF)

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AUTHOR INFORMATION

Corresponding Author

*E-mail: [email protected]. ORCID

Eluvathingal D. Jemmis: 0000-0001-8235-3413 Notes

The authors declare no competing financial interest.

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ACKNOWLEDGMENTS The authors thank IISER-TVM and IISc-Bangalore for computational facilities. E.D.J. thanks the DST for funding through the J. C. Bose Fellowship. J.J. thanks UGC-India for Senior Research Fellowship. I


Article

Inorganic Chemistry

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(14) (a) Riley, K. E.; Hobza, P. Investigations into the Nature of Halogen Bonding Including Symmetry Adapted Perturbation Theory Analyses. J. Chem. Theory Comput. 2008, 4, 232−242. (b) Ř ezác,̌ J.; Riley, K. E.; Hobza, P. Benchmark Calculations of Noncovalent Interactions of Halogenated Molecules. J. Chem. Theory Comput. 2012, 8, 4285−4292. (15) Politzer, P.; Murray, J. S.; Clark, T. σ-Hole Bonding: A Physical Interpretation. In Halogen Bonding I: Impact on Materials Chemistry and Life Sciences; Metrangolo, P., Resnati, G., Eds.; Springer International Publishing, 2015; pp 19−42. (16) Joy, J.; Jose, A.; Jemmis, E. D. Continuum in the X−Z···Y weak bonds: Z = main group elements. J. Comput. Chem. 2016, 37, 270−9. (17) (a) Bauzá, A.; Frontera, A. Aerogen Bonding Interaction: A New Supramolecular Force? Angew. Chem., Int. Ed. 2015, 54, 7340−7343. (b) Bauza, A.; Frontera, A. Theoretical Study on the Dual Behavior of XeO3 and XeF4 toward Aromatic Rings: Lone Pair-pi versus Aerogenpi Interactions. ChemPhysChem 2015, 16, 3625−30. (18) (a) Metrangolo, P.; Murray, J. S.; Pilati, T.; Politzer, P.; Resnati, G.; Terraneo, G. The fluorine atom as a halogen bond donor, viz. a positive site. CrystEngComm 2011, 13, 6593−6596. (b) Metrangolo, P.; Murray, J. S.; Pilati, T.; Politzer, P.; Resnati, G.; Terraneo, G. Fluorine-Centered Halogen Bonding: A Factor in Recognition Phenomena and Reactivity. Cryst. Growth Des. 2011, 11, 4238−4246. (19) (a) Bowmaker, G. A.; Pettinari, C.; Skelton, B. W.; Somers, N.; Vigar, N. A.; White, A. H. Structures and Spectroscopy of some Adducts of Coinage Metal(I) Cyanides with ‘Tetrahedral’-UnidentateN-Bases. Z. Anorg. Allg. Chem. 2007, 633, 415−421. (b) Bowmaker, G. A.; Chaichit, N.; Pakawatchai, C.; Skelton, B. W.; White, A. H. Solventassisted mechanochemical synthesis of metal complexes. Dalton Trans. 2008, 2926−2928. (c) Marcos, C.; Alía, J. M.; Adovasio, V.; Prieto, M.; García-Granda, S. Bis(thiourea)cadmium Halides. Acta Crystallogr., Sect. C: Cryst. Struct. Commun. 1998, 54, 1225−1229. (d) Chieh, C. Synthesis and structure of dichlorobis(thiosemicarbazide)mercury(II). Can. J. Chem. 1977, 55, 1583−1587. (20) (a) Trifan, D. S.; Bacskai, R. Metal-Hydrogen Bonding in Metallocene Compounds. J. Am. Chem. Soc. 1960, 82, 5010−5011. (b) Brammer, L.; McCann, M. C.; Bullock, R. M.; McMullan, R. K.; Sherwood, P. Et3NH+Co(CO)4−: hydrogen-bonded adduct or simple ion pair? Single-crystal neutron diffraction study at 15 K. Organometallics 1992, 11, 2339−2341. (c) Kozelka, J.; Bergès, J.; Attias, R.; Fraitag, J. O−H···PtII: Hydrogen Bond with a Strong Dispersion Component. Angew. Chem., Int. Ed. 2000, 39, 198−201. (d) Aiswaryalakshmi, P.; Mani, D.; Arunan, E. Fe as Hydrogen/Halogen Bond Acceptor in Square Pyramidal Fe(CO)5. Inorg. Chem. 2013, 52, 9153− 9161. (21) (a) Werlé, C.; Bailly, C.; Karmazin-Brelot, L.; Le Goff, X.-F.; Ricard, L.; Djukic, J.-P. Hemichelation, a Way To Stabilize ElectronUnsaturated Complexes: The Case of T-Shaped Pd and Pt Metallacycles. J. Am. Chem. Soc. 2013, 135, 17839−17852. (b) Werlé, C.; Karmazin, L.; Bailly, C.; Ricard, L.; Djukic, J.-P. Stabilization of an Electron-Unsaturated Pd(I)−Pd(I) Unit by Double Hemichelation. Organometallics 2015, 34, 3055−3064. (c) Werlé, C.; Hamdaoui, M.; Bailly, C.; Le Goff, X.-F.; Brelot, L.; Djukic, J.-P. Electron-Deficient η1-Indenyl,η3-allylpalladium(II) Complexes Stabilized by Fluxional Non-covalent Interactions. J. Am. Chem. Soc. 2013, 135, 1715−1718. (22) (a) Einstein, F. W. B.; Jones, T.; Pomeroy, R. K.; Rushman, P. 18-Electron complexes as ligands.Synthesis, structure, and stereochemical nonrigidity of (trimethylphosphine)tetracarbonylosmiumpentacarbonyltungsten complex ((Me3P) (OC)4OsW(CO)5). J. Am. Chem. Soc. 1984, 106, 2707−2708. (b) Fleming, M. M.; Pomeroy, R. K.; Rushman, P. Dissociation and isomerization of (OC)5OsRu(CO)3(SiCl3) (Br), a compound with an osmiumruthenium donoracceptor bond. J. Organomet. Chem. 1984, 273, C33−C38. (c) Grimme, S.; Djukic, J.-P. The Crucial Role of Dispersion in the Cohesion of Nonbridged Binuclear Os → Cr and Os → W Adducts. Inorg. Chem. 2010, 49, 2911−2919. (23) (a) Alvarez, S.; Alemany, P.; Aullon, G.; Palacios, A.; Novoa, J. Heterodox Bonding Effects between Transition Metal Atoms. In The

REFERENCES

(1) (a) Gilli, G.; Gilli, P. The Nature of the Hydrogen Bond: Outline of a Comprehensive Hydrogen Bond Theory; OUP: Oxford, 2009. (b) Hobza, P.; Havlas, Z. Blue-Shifting Hydrogen Bonds. Chem. Rev. 2000, 100, 4253−4264. (c) Scheiner, S. Hydrogen Bonding: A Theoretical Perspective; Oxford University Press: 1997. (d) Joseph, J.; Jemmis, E. D. Red-, Blue-, or No-Shift in Hydrogen Bonds: A Unified Explanation. J. Am. Chem. Soc. 2007, 129, 4620−4632. (2) Grabowski, S. Hydrogen Bonding - New Insights; Springer: Netherlands, 2006. (3) (a) Shahi, A.; Arunan, E. Hydrogen bonding, halogen bonding and lithium bonding: an atoms in molecules and natural bond orbital perspective towards conservation of total bond order, inter- and intramolecular bonding. Phys. Chem. Chem. Phys. 2014, 16, 22935−52. (b) Esrafili, M. D.; Juyban, P.; Solimannejad, M. Exploring lithium bonding interactions between noble-gas hydrides HXeY and LiX molecules (Y = H, CN, NC and X = H, CN, NC, OH, NH2, CH3): A theoretical study. Comput. Theor. Chem. 2014, 1027, 84−90. (4) (a) Zhong, A.; Chen, D.; Li, R. Revisiting the beryllium bonding interactions from energetic and wavefunction perspectives. Chem. Phys. Lett. 2015, 633, 265−272. (b) Villanueva, E. F.; Mo, O.; Yanez, M. On the existence and characteristics of π-beryllium bonds. Phys. Chem. Chem. Phys. 2014, 16, 17531−17536. (5) (a) Grabowski, S. J. Boron and other Triel Lewis Acid Centers: From Hypovalency to Hypervalency. ChemPhysChem 2014, 15, 2985− 2993. (b) Grabowski, S. J. pi-Hole Bonds: Boron and Aluminum Lewis Acid Centers. ChemPhysChem 2015, 16, 1470−9. (6) (a) Bauzá, A.; Mooibroek, T. J.; Frontera, A. Tetrel-Bonding Interaction: Rediscovered Supramolecular Force? Angew. Chem., Int. Ed. 2013, 52, 12317−12321. (b) Mani, D.; Arunan, E. The X−C···Y (X = O/F, Y = O/S/F/Cl/Br/N/P) ’carbon bond’ and hydrophobic interactions. Phys. Chem. Chem. Phys. 2013, 15, 14377−14383. (7) (a) Scheiner, S. The interplay between charge transfer, rehybridization, and atomic charges in the internal geometry of subunits in noncovalent interactions. Int. J. Quantum Chem. 2015, 115, 28−33. (b) Tripathi, G.; Badi-uz-zama, K.; Ramanathan, G. N···N pnicogen bonds in Boc-DOPA-OMe. Chem. Phys. Lett. 2016, 653, 117−121. (8) (a) Fanfrlík, J.; Přad́ a, A.; Padělková, Z.; Pecina, A.; Machácě k, J.; Lepšík, M.; Holub, J.; Růzǐ čka, A.; Hnyk, D.; Hobza, P. The Dominant Role of Chalcogen Bonding in the Crystal Packing of 2D/3D Aromatics. Angew. Chem., Int. Ed. 2014, 53, 10139−10142. (b) Fick, R. J.; Kroner, G. M.; Nepal, B.; Magnani, R.; Horowitz, S.; Houtz, R. L.; Scheiner, S.; Trievel, R. C. Sulfur−Oxygen Chalcogen Bonding Mediates AdoMet Recognition in the Lysine Methyltransferase SET7/9. ACS Chem. Biol. 2016, 11, 748−754. (9) (a) Kolár,̌ M. H.; Hobza, P. Computer Modeling of Halogen Bonds and Other σ-Hole Interactions. Chem. Rev. 2016, 116, 5155− 5187. (b) Wang, C.; Danovich, D.; Mo, Y.; Shaik, S. On The Nature of the Halogen Bond. J. Chem. Theory Comput. 2014, 10, 3726−3737. (10) (a) Bauza, A.; Frontera, A. π-Hole aerogen bonding interactions. Phys. Chem. Chem. Phys. 2015, 17, 24748−24753. (b) Esrafili, M. D.; Asadollahi, S.; Vakili, M. Investigation of substituent effects in aerogenbonding interaction between ZO3 (Z = Kr, Xe) and nitrogen bases. Int. J. Quantum Chem. 2016, 116, 1254−1260. (11) Joy, J.; Jemmis, E. D.; Vidya, K. Negative hyperconjugation and red-, blue- or zero-shift in X−Z···Y complexes. Faraday Discuss. 2015, 177, 33−50. (12) (a) Clark, T.; Hennemann, M.; Murray, J.; Politzer, P. Halogen bonding: the σ-hole. J. Mol. Model. 2007, 13, 291−296. (b) Politzer, P.; Murray, J. S.; Lane, P. σ-Hole bonding and hydrogen bonding: Competitive interactions. Int. J. Quantum Chem. 2007, 107, 3046− 3052. (c) Politzer, P.; Murray, J. S.; Clark, T. Halogen bonding and other σ-hole interactions: a perspective. Phys. Chem. Chem. Phys. 2013, 15, 11178−11189. (13) Ban, Q.; Li, R.; Li, Q.; Li, W.; Cheng, J. Symmetry-adapted perturbation theory interaction energy decomposition for H2CY-XF (Y = O, S, Se; X = H, Li, Cl) complex. Comput. Theor. Chem. 2012, 991, 88−92. J


Article

Inorganic Chemistry Synergy Between Dynamics and Reactivity at Clusters and Surfaces; Farrugia, L., Ed.; Springer: Netherlands, 1995; Vol. 465, pp 241−255. (b) Schwabe, T.; Grimme, S.; Djukic, J.-P. Noncovalent Metal−Metal Interactions: The Crucial Role of London Dispersion in a Bimetallic Indenyl System. J. Am. Chem. Soc. 2009, 131, 14156−14157. (c) Bonifaci, C.; Ceccon, A.; Santi, S.; Mealli, C.; Zoellner, R. W. Cofacial and antarafacial indenyl bimetallic isomers: a descriptive MO picture and implications for the indenyl effect on ligand substitution reactions. Inorg. Chim. Acta 1995, 240, 541−549. (24) (a) Ball, G. E.; Darwish, T. A.; Geftakis, S.; George, M. W.; Lawes, D. J.; Portius, P.; Rourke, J. P. Characterization of an organometallic xenon complex using NMR and IR spectroscopy. Proc. Natl. Acad. Sci. U. S. A. 2005, 102, 1853−1858. (b) McMaster, J.; Portius, P.; Ball, G. E.; Rourke, J. P.; George, M. W. Density Functional Theoretical Studies of the Re−Xe Bonds in Re(Cp) (CO) (PF3)Xe and Re(Cp) (CO)2Xe. Organometallics 2006, 25, 5242−5248. (25) (a) Cobar, E. A.; Khaliullin, R. Z.; Bergman, R. G.; HeadGordon, M. Theoretical study of the rhenium−alkane interaction in transition metal−alkane σ-complexes. Proc. Natl. Acad. Sci. U. S. A. 2007, 104, 6963−6968. (b) Brayshaw, S. K.; Sceats, E. L.; Green, J. C.; Weller, A. S. C−C σ complexes of rhodium. Proc. Natl. Acad. Sci. U. S. A. 2007, 104, 6921−6926. (26) (a) Rodgers, M. T.; Armentrout, P. B. Cationic Noncovalent Interactions: Energetics and Periodic Trends. Chem. Rev. 2016, 116, 5642−87. (b) Nose, H.; Rodgers, M. T. Influence of the d Orbital Occupation on the Structures and Sequential Binding Energies of Pyridine to the Late First-Row Divalent Transition Metal Cations: A DFT Study. J. Phys. Chem. A 2014, 118, 8129−8140. (27) (a) Richardson, M. F.; Wulfsberg, G.; Marlow, R.; Zaghonni, S.; McCorkle, D.; Shadid, K.; Gagliardi, J.; Farris, B. Coordination of ortho chlorines in nickel and zinc 4-substituted 2,6-dichlorophenolates. Crystal and molecular structures of (N,N,N′,N′-tetramethyl-1,2ethanediamine)bis(2,4,6-trichlorophenolato-O,Cl)nickel(II) and tris(pyridine)bis(2,4,6-trichlorophenolato-O)nickel(II). Inorg. Chem. 1993, 32, 1913−1919. (b) Schofield, J. A.; Brennessel, W. W.; Urnezius, E.; Rokhsana, D.; Boshart, M. D.; Juers, D. H.; Holland, P. L.; Machonkin, T. E. Metal−Halogen Secondary Bonding in a 2,5Dichlorohydroquinonate Cobalt(II) Complex: Insight into Substrate Coordination in the Chlorohydroquinone Dioxygenase PcpA. Eur. J. Inorg. Chem. 2015, 2015, 4643−4647. (28) Machonkin, T. E.; Boshart, M. D.; Schofield, J. A.; Rodriguez, M. M.; Grubel, K.; Rokhsana, D.; Brennessel, W. W.; Holland, P. L. Structural and Spectroscopic Characterization of Iron(II), Cobalt(II), and Nickel(II) ortho-Dihalophenolate Complexes: Insights into Metal−Halogen Secondary Bonding. Inorg. Chem. 2014, 53, 9837− 9848. (29) (a) Spackman, M. A.; Jayatilaka, D. Hirshfeld surface analysis. CrystEngComm 2009, 11, 19−32. (b) McKinnon, J. J.; Spackman, M. A.; Mitchell, A. S. Novel tools for visualizing and exploring intermolecular interactions in molecular crystals. Acta Crystallogr., Sect. B: Struct. Sci. 2004, 60, 627−668. (30) Bhattacharya, D.; Maji, S.; Pal, K.; Sarkar, S. Oxygen−Cobalt Chemistry Using a Porphyrinogen Platform. Inorg. Chem. 2009, 48, 6362−6370. (31) Rheingold, A. L.; Liable-Sands, L. M.; Incarvito, C. L.; Trofimenko, S. Novel scorpionate ligands devoid of C-H bonds: BpBr3 and TpBr3. J. Chem. Soc., Dalton Trans. 2002, 2297−2301. (32) Marcos, C.; Alia, J. M.; Adovasio, V.; Prieto, M.; Garcia-Granda, S. Bis(thiourea)cadmium Halides. Acta Crystallogr., Sect. C: Cryst. Struct. Commun. 1998, 54, 1225−1229. (33) (a) Murray, J.; Lane, P.; Politzer, P. Expansion of the σ-hole concept. J. Mol. Model. 2009, 15, 723−729. (b) Politzer, P.; Lane, P.; Concha, M.; Ma, Y.; Murray, J. An overview of halogen bonding. J. Mol. Model. 2007, 13, 305−311. (34) Fraga, S.; Karwowski, J.; Saxena, K. M. S. Handbook of Atomic Data; Elsevier Scientific Publishing Co., 1976. (35) Pyykkö, P. Dirac-Fock One-Centre Calculations Part 8. The 1Σ States of ScH, YH, LaH, AcH, TmH, LuH and LrH. Phys. Scr. 1979, 20, 647.

(36) (a) Buijse, M. A.; Baerends, E. J. Analysis of nondynamical correlation in the metal−ligand bond. Pauli repulsion and orbital localization in MnO−4. J. Chem. Phys. 1990, 93, 4129−4141. (b) Ziegler, T.; Cheng, W.; Baerends, E. J.; Ravenek, W. Theoretical study on the difference in the relative strengths of metal-hydrogen and metal-methyl bonds in complexes of early transition metals and complexes of middle to late transition metals. Inorg. Chem. 1988, 27, 3458−3464. (c) Pyykkö, P. The Physics behind Chemistry and the Periodic Table. Chem. Rev. 2012, 112, 371−384. (d) Jiang, N.; Schwarz, W. H. E.; Li, J. Theoretical Studies on Hexanuclear Oxometalates [M6L19]q− (M = Cr, Mo, W, Sg, Nd, U). Electronic Structures, Oxidation States, Aromaticity, and Stability. Inorg. Chem. 2015, 54, 7171−7180. (37) Pyykko, P. Predicted Chemical Bonds between Rare Gases and Au+. J. Am. Chem. Soc. 1995, 117, 2067−2070. (38) Haynes, W. M. CRC Handbook of Chemistry and Physics, 95th ed.; CRC Press, 2014. (39) Clark, T.; Politzer, P.; Murray, J. S. Correct electrostatic treatment of noncovalent interactions: the importance of polarization. WIREs Comput. Mol. Sci. 2015, 5, 169−177. (40) Politzer, P.; Murray, J. S.; Clark, T. Mathematical modeling and physical reality in noncovalent interactions. J. Mol. Model. 2015, 21, 1− 10. (41) Murray, J. S.; Politzer, P. The electrostatic potential: an overview. WIREs Comput. Mol. Sci. 2011, 1, 153−163. (42) Albright, T. A.; Burdett, J. K.; Whangbo, M.-H. Orbital Interactions in Chemistry; Wiley: New York, 1985; pp xv, 447. (43) 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.; Nakatsuji, H.; Caricato, M.; Li, X.; Hratchian, H. P.; Izmaylov, A. F.; Bloino, J.; Zheng, G.; Sonnenberg, J. L.; Hada, M.; Ehara, M.; Toyota, K.; Fukuda, R.; Hasegawa, J.; Ishida, M.; Nakajima, T.; Honda, Y.; Kitao, O.; Nakai, H.; Vreven, T.; Montgomery, J. A.; Peralta, J. E.; Ogliaro, F.; Bearpark, M.; Heyd, J. J.; Brothers, E. N.; Kudin, K. N.; Staroverov, V. N.; Kobayashi, R.; Normand, J.; Raghavachari, K.; Rendell, A. P.; Burant, J. C.; Iyengar, S. S.; Tomasi, J.; Cossi, M.; Rega, N.; Millam, J. M.; Klene, M.; Knox, J. E.; Cross, J. B.; Bakken, V.; Adamo, C.; Jaramillo, J.; Gomperts, R.; Stratmann, R. E.; Yazyev, O.; Austin, A. J.; Cammi, R.; Pomelli, C.; Ochterski, J. W.; Martin, R. L.; Morokuma, K.; Zakrzewski, V. G.; Voth, G. A.; Salvador, P.; Dannenberg, J. J.; Dapprich, S.; Daniels, A. D.; Farkas, Ö .; Foresman, J. B.; Ortiz, J. V.; Cioslowski, J.; Fox, D. J. Gaussian 09, Revision D.01; Gaussian, Inc.: Wallingford, CT, 2013. (44) (a) Richter-Addo, G. B.; Wheeler, R. A.; Hixson, C. A.; Chen, L.; Khan, M. A.; Ellison, M. K.; Schulz, C. E.; Scheidt, W. R. Unexpected Nitrosyl-Group Bending in Six-Coordinate {M(NO)}6 σBonded Aryl(iron) and -(ruthenium) Porphyrins. J. Am. Chem. Soc. 2001, 123, 6314−6326. (b) Ghosh, A. Metalloporphyrin−NO Bonding: Building Bridges with Organometallic Chemistry. Acc. Chem. Res. 2005, 38, 943−954. (45) (a) Mingos, D. M. P. General bonding model for linear and bent transition metal-nitrosyl complexes. Inorg. Chem. 1973, 12, 1209− 1211. (b) Bohr, F.; Chermette, H.; Ruiz-lópez, M. F. A density functional study of pseudotetrahedral metal−nitrosyl complexes. Int. J. Quantum Chem. 1994, 52, 1039−1049. (46) (a) Pyykkö, P. Strong Closed-Shell Interactions in Inorganic Chemistry. Chem. Rev. 1997, 97, 597−636. (b) Pyykkö, P. Theoretical Chemistry of Gold. Angew. Chem., Int. Ed. 2004, 43, 4412−4456. (47) Dahl, E. W.; Baddour, F. G.; Fiedler, S. R.; Hoffert, W. A.; Shores, M. P.; Yee, G. T.; Djukic, J.-P.; Bacon, J. W.; Rheingold, A. L.; Doerrer, L. H. Antiferromagnetic coupling across a tetrametallic unit through noncovalent interactions. Chem. Sci. 2012, 3, 602−609. (48) (a) Becke, A. D.; Edgecombe, K. E. A simple measure of electron localization in atomic and molecular systems. J. Chem. Phys. 1990, 92, 5397−5403. (b) Silvi, B.; Savin, A. Classification of chemical bonds based on topological analysis of electron localization functions. Nature 1994, 371, 683−686. (49) (a) Broclawik, E.; Stępniewski, A.; Radoń, M. Nitric oxide as a non-innocent ligand in (bio-)inorganic complexes: Spin and electron K


Article

Inorganic Chemistry transfer in FeII-NO bond. J. Inorg. Biochem. 2014, 136, 147−153. (b) Collman, J. P.; Farnham, P.; Dolcetti, G. Intramolecular redox equilibriums of cobalt-nitrosyl complexes. J. Am. Chem. Soc. 1971, 93, 1788−1790. (50) Bader, R. F. W. Atoms in Molecules: A Quantum Theory; Clarendon Press: Oxford, 1990; pp xviii, 438. (51) Koch, U.; Popelier, P. L. A. Characterization of C-H-O Hydrogen Bonds on the Basis of the Charge Density. J. Phys. Chem. 1995, 99, 9747−9754. (52) Hugas, D.; Guillaumes, L.; Duran, M.; Simon, S. Delocalization indices for non-covalent interaction: Hydrogen and DiHydrogen bond. Comput. Theor. Chem. 2012, 998, 113−119. (53) Frenking, G.; Matthias Bickelhaupt, F. The EDA Perspective of Chemical Bonding. In The Chemical Bond; Wiley-VCH Verlag GmbH & Co. KGaA: 2014; pp 121−157. (54) (a) Scheiner, S. Detailed comparison of the pnicogen bond with chalcogen, halogen, and hydrogen bonds. Int. J. Quantum Chem. 2013, 113, 1609−1620. (b) Wolters, L. P.; Bickelhaupt, F. M. Halogen Bonding versus Hydrogen Bonding: A Molecular Orbital Perspective. ChemistryOpen 2012, 1, 96−105. (c) Huber, S. M.; Scanlon, J. D.; Jimenez-Izal, E.; Ugalde, J. M.; Infante, I. On the directionality of halogen bonding. Phys. Chem. Chem. Phys. 2013, 15, 10350−10357. (55) Murray, J. S.; Shields, Z. P.-I.; Seybold, P. G.; Politzer, P. Intuitive and counterintuitive noncovalent interactions of aromatic π regions with the hydrogen and the nitrogen of HCN. J. Comput. Sci. 2015, 10, 209−216. (56) Zhao, Y.; Truhlar, D. G. The M06 suite of density functionals for main group thermochemistry, thermochemical kinetics, noncovalent interactions, excited states, and transition elements: two new functionals and systematic testing of four M06-class functionals and 12 other functionals. Theor. Chem. Acc. 2008, 120, 215−241. (57) Weigend, F.; Ahlrichs, R. Balanced basis sets of split valence, triple zeta valence and quadruple zeta valence quality for H to Rn: Design and assessment of accuracy. Phys. Chem. Chem. Phys. 2005, 7, 3297−3305. (58) Bickelhaupt, F. M.; Baerends, E. J. Kohn-Sham Density Functional Theory: Predicting and Understanding Chemistry. In Reviews in Computational Chemistry; John Wiley & Sons, Inc., 2007; Vol. 15, pp 1−86. (59) te Velde, G.; Bickelhaupt, F. M.; Baerends, E. J.; Fonseca Guerra, C.; van Gisbergen, S. J. A.; Snijders, J. G.; Ziegler, T. Chemistry with ADF. J. Comput. Chem. 2001, 22, 931−967. (60) Frenking, G.; Krapp, A. Unicorns in the world of chemical bonding models. J. Comput. Chem. 2007, 28, 15−24. (61) Mitoraj, M.; Michalak, A. Natural orbitals for chemical valence as descriptors of chemical bonding in transition metal complexes. J. Mol. Model. 2007, 13, 347−355. (62) Mitoraj, M.; Michalak, A. Donor−Acceptor Properties of Ligands from the Natural Orbitals for Chemical Valence. Organometallics 2007, 26, 6576−6580. (63) Keith, T. A. AIMAll, Version 16.01.09; Gristmill Software: Overland Park, KS, 2016. aim.tkgristmill.com. (64) Lu, T.; Chen, F. Multiwfn: A multifunctional wavefunction analyzer. J. Comput. Chem. 2012, 33, 580−592.

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