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Halogen Bonding: Unifying Perspectives on Organic and Inorganic Cases Marina Tawfik, and Kelling J. Donald J. Phys. Chem. A, Just Accepted Manuscript • DOI: 10.1021/jp507879w • Publication Date (Web): 24 Sep 2014 Downloaded from http://pubs.acs.org on September 28, 2014

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The Journal of Physical Chemistry

Halogen Bonding: Unifying Perspectives on Organic and Inorganic Cases Marina Tawfik, and Kelling J. Donald* Department of Chemistry, Gottwald Center for the Sciences, University of Richmond, Richmond, Virginia 23173, United States

Abstract We find for distinct classes of halogen bonded complexes (MF3—X---Y) that the ab initio BSSE corrected binding energies (E) and enthalpies (H) are predicted by functions of the form y = A/rn + C. Here X is a halogen atom, Y is a base, r is the X---Y separation, and A, n, and C are constants. The actual value of n (5.5 < n < 7.0 for E) for each class is determined evidently by the availability of the lone pairs on the base – and is insensitive to M, such that all of the complexes of a given base fall on the same curve for y vs. r. Remarkably, several bases show the same behavior in some cases such that just three curves account for 55 MF3I---Y complexes of eleven bases, where M = C, Si, Ge, Sn, and Pb. Two additional bases, THF and NF3, which form especially strong and weak complexes, respectively, are in classes by themselves. Anomalous modes of halogen bonding are identified; in particular, furan forms sigma-hole complexes via carbons 2 and 3 (through the π-system) in the ring in preference to the oxygen site. These results are in line with experimental observations for furan-dihalogen complexes, and several other small MF3I---Y pairs are proposed in this work for experimental interrogation. Instead of halogen bonding, CF4 tends to form weak sigma-hole bonds to bases via the polarized central carbon atom, and new examples of such pro-dative interactions to carbon in CF4 are identified in this work. We find that GeF3I and SnF3I form I---Y halogen bonds of comparable energies to those formed by the smaller and better studied CF3I. PbF3I forms the strongest halogen bond regardless of the identity of the base; SiF3I consistently forms the weakest link. Keywords: Sigma-Hole, Group 14 Halide, Non-covalent, Interaction Energy, pi-System

*Corresponding Author: Tel.: 804-484-1628; Fax: 804-287-1897; E-mail: [email protected]

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1. Introduction Halogen bonding has been covered under other names for decades. 1 - 2 But the analogy to hydrogen bonding has caused this latest ‘halogen bonding’ appellation to stick. The halogen bond (or X-bond; see Figure 1a) is a stabilizing sub-covalent interaction3 - 7 that is characterized 45

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by an alignment between a localized region of positive electrostatic potential on a halogen atom (X) in a molecule (R-X) (the so-called sigma-hole depicted in blue in Figure 1b)8,9,10 and an electron rich center (Y), such as a Lewis base with a lone pair oriented towards and close enough to X (Figure 1a). The Lewis base, Y, may be an atom or ion, a small molecule like NH3, or a functional group (with a nitrogen or oxygen atom, for instance) in a large molecule or on a surface. The key requirements for a halogen bond are a persistent sigma-hole (-hole) on X and a sufficiently accessible electron rich site nearby. In fact, a large molecule with a polarized halogen atom in its structure and a source of lone pairs appropriately positioned may exhibit intra-molecular halogen bonding. (a)

(b)

R

+x

X

Y

-x

Figure 1: (a) A simple model of halogen bonding, and (b) the computed electrostatic potential (ESP) map on the 0.001 iso-density surface for GeF3I. The ESPs on the surface shown are in the range  x for x = 5.3310-2 au. The region of positive potential (the σ-hole) around the extension of the R—X bond axis is in blue. An optimal halogen bonding interaction requires an alignment of the R—X bond (e.g. F3Ge—I) with an electron rich site on Y (e.g. F3Ge—I----NH3). In the map, the I atom points out from the plane of the page, the F atoms point into the page, and the Ge center is hidden below the I atom. 2


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The region of positive electrostatic potential tends to emerge on X (centered around the pole of X outside the R—X bonding region) in cases where the R fragment (e.g. –CF3 in CF3I) is sufficiently polarizing that the bulk of the electron density in the bonding orbital of X in the R—X bond is coaxed into the bonding region. That allows the nucleus to dominate the potential at X along the bond axis outside R-X bond such that a local positive potential arises on X (see Figure 1). Since the size and the atomic polarizability of X increases significantly going down group 17,11,12 it is not surprising (and it is now well known) that for any series of compounds with the general chemical formula R—X, the size and the strength (the maximum positive electrostatic potential) of the σ-hole at X (for a given R group) varies as I > Br > Cl >> F.13,14 The attention that halogen bonding has attracted in recent years is due in part to a proliferation in the number of experimentally identified instances of these interactions. 15 , 16 Halogen bonding interactions have been implicated, for example, in the folding of molecules, the binding of ligands to metals, and the recognition of hormones by their receptors.17 It is apparent, too, that halogen bonding interactions can be utilized in drug design (with inhibitors forming complexes by halogen bonding),18 in crystal engineering19 (as ordering influences, for instance, in the formation of co-crystals), 20 and in other industries.5,19

Even in the area of crystal

engineering, however, the focus has been persistently on organic systems in which R—X is some kind of halogenated carbon compound. In some cases, the electron rich bonding partner Y (the Lewis base) is also an organic compound with an electron rich center such as N or O – e.g. organic acids, amines, and the like. But organic R groups are not privileged in their ability to induce σ-holes or facilitate halogen bonding. In 2010, our research group raised the question of halogen bonding by R—I = MH3I and MF3I where M = Si, Ge, Sn, and Pb.21 Using NH3 as our base of choice, we found that

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the interaction energy E = E(complex) – [E(R—X)+E(Y)] (where E(R—X) and E(Y) are energies obtained for the isolated minimum energy R—X and Y molecules, respectively) is much more negative for F3M—I---NH3 than it is for H3M—I---NH3, for any M. This outcome was expected, since it was already known, for example, that F3CI forms stronger halogen bonds than H3CI. The surprise was that the ordering of the interaction energies for H3M—I---NH3 as a function of M (i.e. C > Si > Ge > Sn > Pb, which privileges C) were radically different for F3M—I---NH3. In the latter case, the interaction energies vary as Pb >> C  Sn > Ge > Si. We managed to show, therefore, that for a given group in the periodic table (going down group 14, for example), the trend in the size of the sigma-hole induced on X and the energies of the X---Y bonds can vary significantly depending on the identities of the other substituents on the M center. Put another way, replacing H with F in H3M—I---Y strengthens the X-bond in all cases, but much more so in the cases when M = Ge, Sn, and Pb than it is able to when M = C or Si. A significant implication of that result, therefore, is that halogenated Sn and Ge compounds or surfaces can be just as effective as the analogous C cases and even slightly superior in forming halogen bonds. With those observations, therefore, we have started to develop an understanding of the nature of halogen bonding by halides bonded to heavy atoms, including the influence of polarization on the central atoms to which X is bonded on the energetics of bonding. Our initial work in ref. 21 focused on complexes formed by a single base (NH3) as part of a larger study of the tuning of σ-holes. In this paper, we demonstrate that F3PbX consistently forms the strongest X-bonds to common bases compared to its lighter group 14 analogues. Ge and Sn are on par with C in this respect, and even stronger in some cases. Consistently, however, F3SiX forms the weakest link. Moreover, we find a few basic rules for the R—X---Y engagement, including a general mathematical function of the form ΔE = A/rn that governs the relationship between E, which we defined above, and the I---Y interatomic separation, r. The

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values for A and n are roughly independent of R, but rely a lot on the accessibility and availability of the lone pair on the bases. 2. Computational Methods All of the molecular structures discussed in this work have been optimized at the Møller-Plesset (MP2(full)) level of theory. 22 The quantities computed at that level include the optimized geometrical parameters of all of the halogen-bonded complexes, their interaction energies, enthalpies, free energies, and harmonic vibrational frequencies. The cc-pVTZ basis sets23 have been employed for all the elements that we consider in this work that precede tin in the periodic table. For the computationally more demanding cases, M = Sn, Pb, and X = I, we used scalarrelativistic energy-consistent small core Dirac-Fock (MDF) effective-core pseudopotentials (ECPs): with 28e- cores for Sn and I, and a 60e- core for Pb (without the spin-orbit parts but including the scalar relativistic effects) along with the corresponding (MDF cc-pVTZ-pp) basis sets.24,25,26 All of our ab initio calculations have been performed using the Gaussian 09 suite of programs.27 The computed interaction energies (E; see the definition in the introduction section above), enthalpies, and free energies have been corrected for basis set superposition errors using the counterpoise correction procedure28 as implemented in the Gaussian 09 software. Molecular orbital and the electrostatic potential representations have been generated using the GaussView graphics software.29 The Chemcraft software30 has been useful for us as well in analyzing our results and for generating some of the molecular graphics included in this report.

3. Results and Discussion 3.1. Geometry and Stability. Over 100 acid-base combinations or complexes have been considered in this work (Table 1). The series includes three different trifluorohalomethanes and the Si, Ge, Sn, and Pb analogues of trifluoroiodomethane interacting with thirteen different bases 5



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The bases that we consider here are either common reagents in organic or inorganic chemistry like water, ammonia, and tetrahydrofuran (THF) or modifications of simple bases, which enabled us to understand better the influence of substituents on the electron donating properties of the electron rich O or N centers in bases. Table 1: List of halomethanes and heavier (R—X) species, and the Lewis bases (Y) considered in this work. R—X (where R- = F3M-) a

Y

CF4

SiF3I

H2O

H2S

NF3

CF3Cl

GeF3I

CH3OCH3

H2Se

NCl3

CF3Br

SnF3I

THF

NH3

NBr3

CF3I

PbF3I

Furan

N(CH3)3

Pyridine (Py) b

F-pyridine (F-Py) This molecule did not form any halogen bond to the bases considered. Instead, we observed a completely different sigma-hole type interaction between the base and the central C atom, which is polarized by the F atoms. The C center has four sigma-holes – one induced opposite each C—F bond. b Perfluoropyridine (C5F5N). a

In this contribution, we are interested primarily in halogen bonding interactions. But several alternative forms of weak attractive interactions are possible between some of the R—X and Y groups that we consider. One example is (-H---F-) hydrogen bonding between NH3 (or H2O) and CF3X. We do not seek to establish here whether the halogen bonding arrangement is the global minimum energy state among other possible modes of interaction between RX and Y. Indeed, the starting structures for our optimizations were oriented such that the halide expected to form the halogen bond (X = Cl, Br, or I) was oriented directly towards the N or the O center on the base. No geometrical restriction was employed during the optimizations, however, and in some cases alternative complexes have been obtained, in particular between the M—F bond of the MF3X molecules and the E—H bond of the H2E bases, where E = O, S, and Se. The key geometrical parameters obtained at the MP2(full) level of theory for the halogen bonded complexes are listed in Table 2 and in Table S1 in the supporting information. The X---Y distances and M—X---Y bond angles, respectively, are plotted in Figures 2 and 3 as well.31 6


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Table 2. Computed X---Y bond distances for the optimized halogen bonding interactions considered in this work. These results have been obtained at the MP2(full) level of theory with the basis sets mentioned in the methods section. The CF4---Y interaction energy data are not included due to the failure of these systems to produce any evidence of F---Y halogen bonding. CF3Cl

CF3Br

CF3I

SiF3I

GeF3I

SnF3I

PbF3I

H 2O

2.998

2.964

3.001

3.174

3.070

3.057

2.916

H 2S

3.581

3.531

3.534

3.820

3.675

3.635

3.349

H2Se

3.652

3.586

3.579

3.882

3.731

3.681

3.355

THF

2.882 (3.221) 3.291 2.907

2.831 (3.239) 3.268 2.857

2.852 (3.375) 3.352 2.889

3.063 (3.421) 3.606 3.103

2.933 (3.364) 3.464 2.983

2.909 (3.346) 3.419 2.958

2.707 (3.251) 3.159 2.766

Pyridine

2.966

2.865

2.850

3.132

2.981

2.928

2.595

F-pyridine

3.100

3.063

3.126

3.324

3.215

3.197

3.013

NH3

3.048

2.970

2.967

3.246

3.099

3.060

2.787

N(CH3)3

2.867

2.751

2.764

3.046

2.889

2.822

2.520

NF3

3.247

3.264

3.351

3.568

3.468

3.465

3.328

NCl3

2.998

2.953

3.014

3.194

3.102

3.083

2.916

Furan* CH3OCH3

NBr3 2.903 2.865 2.926 3.096 3.006 2.981 2.804 THF: tetrahydrofuran; F-Pyridine: perfluoropyridine; *Distances between X and the center of the furan ring are in parentheses. The other value is the shortest X---C distance in the R—X---furan complex.

The relatively short X---Y distances in Figure 2 (which are consistently lower than the sum of the van der Waals radii of X and N, O, S or Se) and the generally very small deviations of the M—X---Y bond angles from 180o (Figure 3) are telltale signs of halogen bonding interactions.

Figure 2: X--(N, O, S, or Se) separations in the F3M—X---Y complexes for X = I. The additional cases for CF3-Cl and CF3-Br are included as well for comparison. Furan is excluded. 7



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The sums of the van der Waals (vdW) radii for X = Cl (1.81 Å), Br (1.95 Å), I (2.15 Å), C (1.85 Å), N (1.54 Å), O (1.40 Å), S (1.85 Å), and Se (2.00 Å) for the pairs of atoms directly involved in the X---Y X-bonds are consistently larger than the X---Y separations. For I and Se the sum of the vdW radii is 4.15, for instance, and the I---Se separations in F3M—I---SeH2 run from 3.882 Å to 3.355 Å (see the H2Se row in Table 2). For Cl---O, the sum of the vdW radii is 3.21 Å but the longest Cl---O separation is 2.998 for Y = H2O. As we see presently, furan does not bond via the oxygen center. The halogen bond is between the halide and the -system at one of the two C—C bonds that is on either side of the O atom in the ring.

Figure 3: M-X--(N, O, S, or Se) bond angles in the F3M—X---Y complexes for X = I, and for CF3-Cl and CF3-Br. The furan cases are excluded. In the F3M-X—furan complexes, the M—X bond is directed towards a C-C bond in the ring (see Figure 4), not towards oxygen. The roughly linear M—X---Y bond angles observed in most of the optimized complexes (Figure 3) are expected for strong halogen bonding interactions where steric hindrance or other competing interactions such as hydrogen bonding are absent (see Figure 4a-c). But we also observed some deviations: Figure 4 is an illustrative summary that includes an alternative form of interaction (4d), other examples of simple X-bonds (4e-g) and exceptions (4h-j); and each case shown represents a minimum on the relevant potential energy surface. The M—X---Y bond 8


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angles deviate noticeably in some cases from 180o – by as much as 10o at the extreme for THF, and to lesser extents from dimethyl ether, and pyridine (Figure 3). The weakest of the NF3 complexes (which we identify on the right in Figure 3) are exceptional. The X---N contacts in those three compounds are relatively long (Figure 2), and the bond angles deviate even more strongly from 180o (see Figure 4h). We examine that and other anomalies in the next sections. Side-view 1

Side-view 2

(b) (c) (d) (a ) (a)Trig.-planar O center in MF3-I---OH2 as observed in this work for M = Si, Ge, and Sn. (bc) Trig.-pyramidal O center in MF3-I---OH2 for M = C, and Pb, respectively. (d) Alternative CF3X--Y interaction observed for X = Cl, Br, I and Y = H2O, H2S and H2Se. For H2S and H2Se, however, one of the H atoms points downwards, away from the CF3X molecule.

(f) MF3-I---(F-Py) (M = Ge)

(g) (h) Tetrahedral I---NR'3 orientation (g) observed in most cases (with an example shown for M = Ge and R = F) and the See-saw type (h) observed for CF3Cl---NR'3 for R = F, and Cl and SiF3I---NF3.

(i) Interaction between X and the furan ring (E.g.: M = C; X = Cl)

(e) MF3-I---THF (M = Ge)

(j) Interaction between X and the furan ring (for M = Ge; X = I)

Figure 4: Representative minimum energy halogen bonded and other structures that we obtained. The MF3-I---OH2 panel (top (a-d)) includes an alternative CF3I---H2O complex that we obtained ‘(d)’, and similar CF3X complexes with H2S and H2Se were located as well (see the supporting information). In almost every case, however, the X-bonded form (b in this figure) is lower in energy. 3.2 Observations on Exceptions 3.2.1: How Water Bonds: The geometry of the F3M—I---OH2 complex depends on the identity of M. The O atom bonds to I in SiF3I, GeF3I, and SnF3I in a roughly trigonal planar fashion (Figure 4a) in which the M—I bond and the H2O molecule are almost coplanar. The computed 9



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‘H-I-O-H’ dihedral angles are 169o, 180o, and 178o for M = Si, Ge, and Sn, respectively, where the coplanar extreme is 180o. For CF3—I---OH2, and PbF3—I---OH2, however, the corresponding angles are 120o and 138o respectively, which are much more in line with the ~120o that would be expected if – following a simple bonding picture – the iodine is pointing directly towards one of two ~sp3 lone pair hybrid orbitals on oxygen. But why would the I---O bonding pattern change going from M = C to Si and again from M = Sn to Pb, especially since the initial geometries for the optimizations were similar, with the M—I bond roughly perpendicular to the H2O molecule? (a)

I---O interactions (for M = Sn) I---O Bonding I---O Antibonding H 2O

I--O bonding region

MF3I

(b)

HOMO-10 HOMO-7 HOMO-2 I---O Interactions (for M = Pb) I---O Bonding I---O Antibonding

HOMO-13 HOMO-12 HOMO-3

HOMO-9 HOMO-2

Figure 5: Molecular orbitals for two F3M—I---OH2 complexes: one case with the trigonal planar O center (M = Sn (top)) and another with the pyramidal arrangement (M = Pb (bottom)). We find, from a molecular orbital perspective, that the I---O interactions can be understood as an overlap between a σ* orbital concentrated on I in MF3I and non-bonding (lone pair) orbital(s) on O in H2O. For M = Si, Ge, and Sn, that picture works very well. The molecular orbitals (MOs) for SnF3I---OH2 in Figure 5a show direct interactions (two bonding and one antibonding) between the I atom (purple) and the in-plane O p orbital of H2O, which favors the 10


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trigonal planar I---OH2 arrangement. The other oxygen lone pair is perpendicular to the I---O bond axis and is uninvolved in the bonding. For M = C, and Pb, however, both lone pairs are involved. The MOs in Figure 5b show that the iodine atom has an indirect but still substantial bonding interaction with both of the oxygen lone pairs (compare HOMO-12 and HOMO-3 in Figure 5b, for example) – rather than a direct interaction with only one. We surmise that the far larger and stronger σ-holes on I when M = C and Pb make it worthwhile energetically to sacrifice a direct interaction with one electron pair on the oxygen center for weaker indirect interactions with both lone pairs that, in sum, produces stronger binding overall. Indeed, only the pyramidal option (Figure 4b,c) is located for M = C or M = Pb. For M = Si, Ge, and Sn, we know from ref. 21 that the σ-hole on I is weaker and smaller than it is for M = C and Pb – so much so it appears that the sum of two indirect interactions is weaker in those cases than a direct interaction with just one of the lone pairs on O. 3.2.1: H---F Hydrogen Bonds Coexisting with other Weak Interactions: The alternative complex that we find for water, H2S, and H2Se (Figure 4d) involves a rough alignment between the polar M(+)—F(-) bond and the polar E(-)—H(+) bond of the H2E molecule. Apart from Figure 4d, graphical representations of these complexes are shown in Figure S1 in the supporting information. Fluorine is the most polarizing of the halides, and it tends to induce massive and strong σ-holes on central atoms to which it is bonded, which is the M center in this case (see refs. 32 - 38). The systems shown in Figures 4d and S1 appear to combine a F---H hydrogen bond and 3334

353637

a weak M---E sigma-hole interaction38 in a way that rivals in strength the simple halogen bonding I---O interaction. These more involved dipole-dipole type interactions that we find in this work, which are generally weaker than the X-bonded forms for M = C (Figure S1), are not considered further in this work.

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3.2.2: Role Reversal - σ-holes on Lewis Bases: The F3M—X---NR3 complexes typically arrange so that the M—X---N contact forms an almost straight line (with bond angles less than 179o), the two fragments fall along a common principal axis, and the N center is pseudotetrahedral (Figure 4g). For F3C—Cl---Y (where Y = NF3, and NCl3) and F3Si—I---NF3, however, the F3M—X and the X---Y axes are not aligned. The Y = NR3 fragment is tilted relative to the F3M—X axis (see Figure 3)39 so that the arrangement of the X---N bond and three R—N bonds around N is closer to a see-saw shape (Figure 4h) than the tetrahedral geometry (Figure 4g). This anomaly may be explained by the relative weakness – already well established (in ref. 21 and references therein) – of halogen bonds when X = Cl, or where X is bonded to silicon.21 But the R—X unit does not bare all the blame for the ill-shaped R—X---NF3 complexes. The halogen bonds formed by NF3 are all quite weak; and that is hardly surprising. After all, the N centers in NF3 (and NCl3) are positively charged. And a consequence of this significant polarization of N by the F (or Cl) substituents is that the lone pair on the NF3 (and to a lesser extent for NCl3) is far less available for bonding than it is in bases – such as N(CH3)3 – where the substituents are less electron withdrawing. In fact, we showed in ref. 38 – a study of dative bonds supported by sigma-hole interactions – that the F4M---NF3 interaction energies are relatively small and the M---N contacts are long for that very reason: F4SiNF3, for example, has an Si---N bond distance of 3.179 Å compared to 2.078 Å for SiF4N(CH3)3.38 So, the NF3 interactions are weak, but why does that foster the observed distortions away from a simple halogen bond? The potential energy surface of F3C—Cl---NF3 shown in Figure 6 uncovers significant σ-holes on both the terminal Cl center (as expected) and on the N atom in NF3 as well! The NF3 unit in the atypical MF3X---NF3 complexes is oriented, therefore, in such 12


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a way that the lone pair (in green on N in ‘side-view 1’ in Figure 6; see also region a of ‘view 2’) interacts with the σ-hole on X, while also allowing the lone pairs (the electron rich belt) on X (see Cl in ‘side-view 1’ and region b in ‘side-view 2’) to interact with one of the three σ-holes on N in NF3. F3CCl---NF3 Side-View 1

Side-View 2 (and vertical)

electron rich Cl belt CF3Cl

b σ-hole

a NH3

In these two views of the same complex, the -holes are obvious on both the CF3Cl and NH3 molecules. The N atom interacts with the -hole on Cl (a) but the CF3Cl tilts a bit so that the electron rich belt (the lone pairs) on Cl can interact as well with one of the -holes on N (b) (see view on the right).

Figure 6: Different modes of σ-hole bonding coexisting. Obvious sigma-holes (blue regions in the figure) exist on both the CF3Cl and the NF3 potential surfaces. The surface was selected to show the -holes and the most significant contacts between the two molecules. The combination of those two compromised interactions appears to be more stabilizing for the complex than either of the direct interactions by themselves. If the σ-holes induced by F3M- on X are sufficiently strong, however, as in the cases where X = Br or I, (or if the electron donor site on the base has no significant σ-hole) distortions are less likely. 3.2.3: The Furan Exception: In line with earlier predictions 40 and experimental microwave data41 -43 for simple furan---dihalogen (and ethene---X—R) complexes, we find that furan forms 42

no halogen bond via the oxygen center. The MF3I species interact instead with the -system at the C(2)-C(3) bond in furan. The bonding MOs in Figure 7a-c show that the orbital interactions that support bonding between furan and the halogen center involve (i) at least one of the perpendicular p orbitals on the iodine center (HOMO-1 and -3), and (ii) the M-I anti-bonding 13



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orbital (which is usually implicated in halogen bonding (HOMO-4 in Figure 7c)). In contrast to the furan case, the pyridine halogen bonded complexes involve a -type interaction with the lone pair on the N center (Figure 4f), which is again in accordance with predictions for the dihalogen complex.44 The N in pyridine is a better  donor for X-bonding than the O site in furan.

(a) HOMO-1 (b) HOMO-3 (c) HOMO-4 Figure 7: MOs for F3C—I---furan. The grey bond linking the red O center and I is useful for orientational purposes only. The I is closer to the adjacent C-C bond (see Figure 4i,j). 3.3 The Strength of the R—X---Y Halogen Bonds: The strength of a given X-bond is sensitive to both the nature of R and the chemical composition of Y well beyond the particular atomic site that is directly involved in the bond. The binding energy of the H3M—I---NH3 complex is strongest when M = C, for instance. But, we showed recently21 that if we replace H with F to give F3M—I---NH3, the binding energies increase regardless of the identity of M, but in such a disproportionate manner that the Pb complex is more strongly bound than all the others, and the binding energies when M = Sn and Ge are actually comparable to those for the CF3I complex. In ref. 21, we considered only one base: NH3. We report herein that the superiority of the PbF3I system compared to the corresponding halomethane (CF3I) is quite general, and that GeF3and SnF3- are typically just as effective as CF3- in inducing σ-holes and fostering halogen bonds to terminal Cl, Br, and I atoms. The computed (MP2(full)) interaction energies before and after basis set superposition energy correction, are listed in Tables 3 (for M = C) and 4 (for all M for

14


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X = I). We have separated the CF3X and the MF3I systems to allow for a straightforward assessment of the X and M dependence of the halogen bond energies. In each case, the isolated R—X and Y molecules were optimized individually and the sum of their energies was subtracted from the total energy of the optimized X-bonded complex. We exclude the CF3—F---Y case (i.e. where X = F) from Table 3, since those complexes showed no hint of halogen bonding.

Table 3. Computed X---Y interaction energies (in kcalmol-1 units) for all of the optimized CF3X---Y bonding interactions. The BSSE corrected interaction energies that are lower than an arbitrary cut-off of -3.5 kcalmol-1 are in bold type. The CF4 case is excluded. We found no evidence of halogen bonding by that species with the bases we considered. Uncorrected Y

BSSE Corrected

CF3Cl

CF3Br

CF3I

CF3Cl

CF3Br

CF3I

H 2O

-2.43

-3.49

-4.61

-1.46

-2.05

-3.04

H 2S

-1.42

-2.07

-2.95

-1.01

-1.47

-2.21

H2Se

-1.61

-2.36

-3.35

-0.95

-1.42

-2.21

THF

-8.90

-10.70

-12.09

-7.39

-8.22

-9.80

Furan

-2.92

-4.08

-4.99

-1.91

-2.39

-3.31

CH3OCH3

-3.34

-4.81

-6.16

-2.05

-2.80

-4.14

Pyridine

-3.66

-5.64

-8.12

-2.59

-3.81

-6.02

F-pyridine

-2.45

-3.51

-4.52

-1.53

-2.05

-2.90

NH3

-2.94

-4.56

-6.60

-2.17

-3.16

-4.82

N(CH3)3

-4.63

-7.52

-10.53

-2.96

-4.54

-7.35

NF3

-0.97

-1.27

-1.49

-0.41

-0.57

-0.74

NCl3

-2.48

-3.53

-4.55

-1.55

-2.07

-2.90

NBr3

-3.37

-4.82

-6.17

-1.90

-2.50

-3.58

For the MF4 molecules in general, the bases tend to align with one of the -holes induced on M by the F atoms to form a bond to M. Those F4M---Y complexes can be quite stable, and they tend to progress from weak pairs when M = C to dative bonds for Si, Ge, Sn, and Pb.38 Among the MF3I complexes, the magnitudes of the computed interaction energies in Table 4 are smallest in each case for the Si systems. Indeed, even if we include the CF3Cl and CF3Br cases that are expected to be relatively weak, the species with the two weakest halogen 15



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bonds (the smallest BSSE corrected energies in Tables 3 and 4) are SiF3I---NF3 and CF3Cl---NF3. Our results confirm unequivocally as well the very well-known trend of an increase in the halogen bond strength going from X = Cl  Br  I for any M. And the strengthening that occurs going from Br  I is, we find, almost always larger than it is from Cl  Br (Table 3). One implication of this result is that – if we were able to trap an instance of it – the R—At molecule may form super-strong halogen bonds, though they would not be particularly relevant for applications given the instability of At. The strength of the halogen bond is influenced directly by both the polarizing power of R, and the softness of X. Yet, we find that certain fundamental aspects of the halogen bonding interaction is dictated by the nature base - not simply the identity of the atom donating the lone pair (be it N, O, or otherwise), but the specific environment in which that atom exists. So, for a given R—X bonding partner, the oxygen centers in two different bases, for example, may form halogen bonds that have radically different interaction energies. Table 4. BSSE corrected and uncorrected X---Y interaction energies (in kcalmol-1 units) for all of the optimized MF3I---Y halogen bonding interactions considered in this work. The lowest corrected interaction energies (lower than an arbitrary cut-off of -3.50 kcalmol-1) are in bold. Y H2O H2S H2Se THF Furan CH3OCH3 Pyridine F-pyridine NH3 N(CH3)3 NF3 NCl3 NBr3

CF3I -4.61 -2.95 -3.35 -12.09 -4.99 -6.16 -8.12 -4.52 -6.60 -10.53 -1.49 -4.55 -6.17

Uncorrected SiF3I GeF3I -3.26 -4.42 -1.90 -2.61 -2.12 -2.90 -10.20 -11.59 -4.28 -5.17 -4.38 -5.64 -5.27 -7.28 -3.34 -4.34 -4.13 -5.83 -6.52 -8.93 -1.04 -1.25 -3.39 -4.24 -4.72 -5.82

SnF3I -4.63 -2.77 -3.12 -11.85 -5.39 -5.85 -7.85 -4.52 -6.28 -9.78 -1.24 -4.41 -6.11

PbF3I -6.22 -4.42 -5.28 -14.55 -7.17 -8.16 -12.96 -6.16 -9.85 -17.20 -1.61 -5.97 -8.26

CF3I -3.04 -2.21 -2.21 -9.80 -3.31 -4.14 -6.02 -2.90 -4.82 -7.35 -0.74 -2.90 -3.58

BSSE Corrected SiF3I GeF3I SnF3I -2.20 -3.17 -3.39 -1.40 -1.97 -2.15 -1.38 -1.94 -2.15 -8.27 -9.36 -9.70 -2.84 -3.46 -3.72 -2.74 -3.73 -3.99 -3.71 -5.33 -5.90 -2.08 -2.80 -3.03 -2.91 -4.30 -4.69 -4.17 -5.99 -6.75 -0.48 -0.59 -0.61 -2.14 -2.71 -2.91 -2.69 -3.40 -3.69

PbF3I -4.66 -3.50 -3.73 -12.04 -5.11 -5.92 -10.18 -4.34 -7.66 -13.02 -0.85 -4.08 -5.23

16


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CF3I---THF and CF3I---OH2 have binding energies, for instance, of -9.80 and -3.04 kcalmol-1, respectively, even though O is the electron donor in both cases. And similar instances are readily identified in Tables 3 and 4 for nitrogen bases. In general, halogenation of the nitrogen center (and, more broadly, substituting for electron-withdrawing group on the base) weakens the halogen bond. As we mentioned, NF3 forms the weakest contacts of all the bases – the corrected interaction energies are never better than -1.0 kcalmol (Tables 3, 4, and S2), and the corresponding enthalpies and free energies (Tables S3 and S4) are all positive. Trimethylamine and pyridine form, however, some of the strongest bonds in the series (Table 4). 3.4 Length and Strength: We probed further the relationship between the X---base distances and the strengths of the halogen bonds. Shorter halogen bonds are indicative typically of stronger halogen bonds (if the two atoms directly involved in the bond, I---N, for example, are the same in the complexes being compared). We wanted to understand better, however, the relationship between interaction energies and X---base separations and how that relationship reflects fundamental aspects of the nature of the halogen bond. To date, for instance, we have no quantitative picture of distance dependence of halogen bonds as functions of R substituents (inorganic or organic) or different classes of electron donors spanning different rows of the periodic table. The coordinates of the halogen bonded complexes considered in this work are included in the supporting information in Table S5. Figure 8 is a plot of BSSE corrected interaction energies vs. the I---base bond distances for the systems considered in this work (Tables 1, 2, and 4). With the significant differences among the bases that we considered and the orbital effects that are naturally added and altered going from C to Pb down group 14, we did not expect any simple

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relationship between the lengths and strengths of the halogen bonding interactions. But we looked nonetheless, and the patterns that emerge (Figure 8) are remarkable. The sixty-five data points of the corrected interaction energies, E, (and enthalpies, H) vs. the corresponding X-bond distances for the iodides separate neatly into five subsets or series (Figure 8). In the three graphs, series 1 includes, for example (see the caption to Figure 8), all of the complexes of NCl3, NBr3, H2O, and O(CH3)2. An outstanding feature of each of these series is that they are all independent of the identity of R. Moreover, they show a simple inverse (A/rn + C) relationship between E (and H) and the X-bond distance. The constant A varies significantly from one series to another for E (or H) and n is within the range 5.8 – 6.7 for E (and 3.2 to 4.4 for H). We set C = 0 for E imposing the condition that the R—I---Y interaction energy goes to zero as the I---Y distance tends to infinity.

18


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(a)

(b)

(c)

Figure 8(a-c): Corrected interaction energies, enthalpies, and free energies (in kcalmol-1 units) for F3MI---Y X-bonding interactions. Y = Series 1: NCl3, NBr3, H2O, and O(CH3)2, Series 2: NH3, N(CH3)3, pyridine, and perfluoropyridine (F-pyridine), Series 3: H2S and H2Se, and furan (). The NF3(), and THF () systems appear to follow separate curves and are treated as independent cases. Different shapes are used in series 3 (in blue) to distinguish the H2E molecules from furan (). Furan bonds in an unusual way (via the C atoms in the ring) but falls well into series 3. 19



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There are two other aspects of interest from Figure 8. (i) The graph of E (or H) vs. r cannot be readily partitioned into clear categories using either the identity of central atom M in R—X only, or simply the identity of the donor atom in Y (be it N, O, S, or Se). For any given base (Y), all of the MF3I---Y complexes fall on the same curve (i.e. they are in the same series) in Figure 8, but the complexes of bases that have the same donor atom (e.g., the N bases) do not all fall into the same series. N is in the NF3 series, series 1, and series 2, for instance, and O is in THF series, and series 1. Additionally, (ii) the free energies (G) show no ordering of the sort that we observe for E and H. This situation reflects obvious differences in the structure and composition of the bases and the different extents to which entropic effects influence the stability of the complexes. The absence of any similar trend in the G data is not surprising, therefore, given the significant differences in the number of atoms in and the geometries of the various bases that are considered in this work. How, though, do we rationalize the patterns in E and H vs. r, especially the fact that the base is the ordering factor, independent of R- in R—I? The nature of the interaction is determined by I and Y (and not simply the identity of the donor atom in Y but the entire chemical environment in which the N or O atom is situated). The charge density on those electron donor sites and the accessibility of their lone pair(s) is controlled by the substituents. The R- fragment in the partner R—I molecule tunes the strength of the X-bond and determines where the complex ends up along a given curve in Figure 8a,b, but Y dictates the curve on which the complex falls. The ordering that we have identified in Figure 8 emerged once the data were plotted and the complexes were identified. But an even finer categorization may yet be identified once more bases are studied. One could already consider, for instance, that furan (blue triangles in Figure 8), which shows some deviation from the best-fit curve for series 3, belongs to a separate series. 20


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We hope to have shown, however, that a classification of halogen bonds according to bases is possible. A similar ordering that relies solely on the identity of R will probably not be achieved. 3.5 A Test for Universality: Chlorine and bromine are smaller and less polarizable than iodine. So, compared to the I---Y data in Figure 8, the corresponding interaction energies (enthalpies and free energies) would appear to the right of and higher up in energy (since Cl---Y and Br---Y are weaker than the I---Y X-bonds for any given R and Y). To test the universality of the partitioning achieved in Figure 8, we have added the F3C—Cl--Y and the F3C—Br—Y data (Tables 2 and 3) to the graph of the interaction energies in Figure 8a (see Figure 9). To level the playing field, however, (given the size differences between Cl, Br, and I) we multiplied each X---Y bond distance by the following ad hoc scaling factor:  = van der Waals radius of iodine / van der Waals radius of X

Equation 1.

The van der Waals radii used are those mentioned in section 3.1, such that the  = 1.188, 1.103, and 1.000 for Cl, Br, and I respectively.

Figure 9: Corrected interaction energies, (in kcalmol-1 units) for F3MI---Y, F3MBr---Y and F3MCl---Y X-bonds as functions of the scaled X---Y bond distances. The Cl and Br data points are colored according to the series in which the bases belong (as defined above), but they are left unshaded to allow for comparison and an assessment by the reader of how well they fit. NOTE: The best-fit curves from Figure 8 have been extrapolated, but not refitted. 21



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This scaling factor is simple, but it achieves a gratifying collapse of the Cl and Br data well in line with the trends that we observed for the iodides alone in Figure 8. The energies of the Cl and Br complexes (the unshaded data points) are partitioned well (though not perfectly) into the five series that we identified in Figure 8. Moreover, the NF3 and the THF values fall precisely in line with those exceptional data sets (in black in Figure 8), thus providing substantial support for the conclusion that these two bases (NF3 and THF) belong to two separate series for which our set of bases include no other example. The scaling extends the utility of the curves for predicting energies for any halide bonded to a given base. So, at least five sets of base profiles have emerged. THF forms especially strong halogen bonds, despite any on-site repulsive interaction between the R—X units and the terminal hydrogens on the THF ring. Series 3 has relatively long bond lengths because of the sizes of the atomic sites involved (S, Se, and the expansive π system in furan). The shift to series 2 and 1 occurs as the donor atoms get smaller (going to the electron rich and negative nitrogen centers in series 2 to the oxygen and polarized nitrogen centers in series 1, which includes NCl3 and NBr3). The most polarized case that we consider – NF3 – is therefore, perhaps unsurprisingly now, in a class by itself. To sum up, we make the following basic observations: (i) As R becomes more electron-withdrawing, the -hole on X in R—X intensifies, the X---Y bond strengthens and we slide down the curves in Figure 9, but we stay on the same curve for any given Y. (ii) The strength of the bond to a given electron donor site (N or E in Y = NR3 or ER2, for example) decreases and the X---Y distance tends to increase (a shift up the curve) as R becomes more electron-withdrawing. However, (iii) A change in R can also lead to a shift out of the series entirely. Sufficiently bulky or electronegative R substituents can modify the bonding environment at the electron rich site in the base so dramatically as to modify the character of the X-bond interaction drastically. This shows up as a shift from one curve to another in Figures 8 and 9. - For instance, the N center is negative in N(CH3)3 in series 2 – as might be expected for a base – but that same ‘electron rich’ N center has a net positive charge in NF3, and the overall interaction of that site with the positive -hole is weakened accordingly. 22


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3.6. Alternative Bonding Mode for MF4: Halogen bonds to terminal F atoms are rare. For the systems we consider in this work, not even one case of any F---Y halogen bonding interaction has been identified. The reason for this is readily explained by the substantial electron withdrawing character of F, which typically enforces an overall net negative electrostatic potential over the entire surface of the F atom in an R—F molecule. R would have to be extremely electron withdrawing itself to stabilize a significant σ-hole on F, and the fluorinated group 14 centers we consider in this work fail in that effort. But that is only half the story; as we mentioned above, F polarizes the M center to which it is bonded, and can induce substantial σholes that can interact (just as sigma-holes on terminal halides do) with bases as well. CF4·NBr3

CF4·N(CH3)3

CF4·THF

CF4·F-Pyridine

CF4·Pyridine

Figure 10: The many ways to have a CF4Base bond. The examples shown are complexes of CF4 with NBr3, N(CH3)3, THF, F-pyridine, and pyridine. The electrostatic potential on the 0.01 au isodensity surface for each complex is shown in the second row. We have chosen a rather low surface to look at only because it gives us an uncluttered view of the interactions in the bonding region, exposing more clearly the σ-holes (blue) on C (and, yes, also on some terminal atoms such as Br in NBr3 and H in pyridine) and the location of N or O lone pairs (in yellow) on the base.

These sigma-holes arise on M opposite each M—F bond (blue regions on the C centers in Figure 10), which is exposed in the tetrahedral MF4 molecule. So, in MF4 a σ-hole is present in all four basins on the molecular surface, each surrounded by three F atoms (and opposite the fourth F 23



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atom that induced the σ-hole). The complexes formed by the tetrahalides are quite different, therefore, from the halogen bonded species. In general, the base attacks at the σ-hole opposite one of the M—F bonds to form a complex with a locally trigonal bipyramidal geometry at the M center (see, for example, the CF4NBr3 complexes in Figure 10). We examined recently a series of such MF4NR3 complexes for all M,38 and others have found experimental and computational cases for M = C, Si, and Ge;32-37,45, 46 see, also, references 18-20, 37, and 38 in reference 37. This kind of σ-hole supported MF4Y bonding can be quite strong for M = Si, Ge, Sn, and Pb, where empty valence M orbitals coincide with the location of the σ-hole.38 And examples of similar dative type bonds to groups 15 and 16 atoms have been identified and even named (pnictogen and chalcogen bonding) as well. We prefer to avoid naming these bonds to central atoms in the way we do for hydrogen or halogen bonds since, for elements below period 2 (under C, N, and O in groups 14, 15, and 16) hypervalence and strong dative bonding occur commonly. These covalent bonds may be fortified by a σ-hole that coincides with the bond, but they are not weak primarily σ-hole interactions as the ‘element’-bonding eponym might suggest. CF4NH3, for example, has a long and weak C---N bond of 3.206 Å, while the corresponding Si---N, and Pb--N bonds are far shorter and stronger at 2.078 Å, and 2.293 Å, respectively.38 Alternative interactions observed for CF4 are described in reference 38. For H2O, 47,48 the complex is similar in arrangement to that shown in Figure 4d, and we have found more recently similar cases of an alignment of the C—F and E—H bonds for E = S and Se. For both furan and F-pyridine, which we did not consider in the ref. 38, we find an unusual situation in which the base interacts with the σ-hole on C in CF4 via the C(3) atom two bonds away (see the CF4pyridine case in Figure 10), and not via the O or the N atom in the ring. This observation is in line with what we find in the MF3I---furan complex that we reported above (Figure 4i,j) where the electron donor is the C(2)—C(3) bond in the ring, not the O center.

24


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4. Summary and Outlook The distance dependence of covalent bond energies across groups of molecules has been of interest for a long time. More than half a century ago, for instance, Pauling identified a 1/r dependence of the dissociation energies of homonuclear diatomic molecules (A2) with internuclear separation (r) where all the A atoms are in the same group of the periodic table.49 We show in this report that the interaction energies of the halogen bonds – far weaker than classical covalent bonds – exhibit a distance dependence of the form 1/rn, where n is between 5.8 and 6.7. The exact value of n depends on the identity of the base in the complex. A single value of n unites all the R—I---Y complexes for each base in a way that is surprisingly independent of the identity of R (Figure 8). Several bases show essentially the same relationship between E (or H) and r so that just three curves, each with R2 > 0.91, account for the complexes of eleven different bases (Series 1 – 3 in Figure 8). Moreover, after scaling for differences in the sizes of Cl and Br atoms, the interaction energies for the Cl---Y and Br---Y halogen bonds fall rather neatly into the corresponding I---Y series (Figure 9). Apart from simple linear (F3M—I---Y) halogen bonds, various unusual alternative modes of interaction have been identified. In the furan complexes, the C centers in the ring are a better electron donors for halogen bonding than the O center, specifically the C(2)—C(3) (or C(4)— C(5)) bond in the ring. Alternative modes of interaction such as a dipole-dipole interaction that combines ostensibly a sigma-hole interaction between M and E, and a hydrogen bond between F and H are identified for the F3M—I---H2E complexes where E = O, S, and Se. The thermodynamics of the halogen bonding interactions have been assessed. Tetrahydrofuran and trimethylamine form the strongest complexes for any given ‘acid’ (F3M—I) partner. Among the full set of F3M—I systems considered – all of the cases for the group 14 M

25



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atoms from carbon down the group to lead – it is the Pb species (F3Pb—I) that forms the strongest complexes. This result confirms an observation made initially by our group a few years ago21 that halides can form far stronger halogen bonds when M = Pb compared to the betterstudied case for M = C, if the other substituents on the M center are very electron withdrawing. And, for that same reason, F3Ge—I and F3Sn—I form halogen bonds comparable in strength to those of the F3C—I complexes. Strong halogen bonding interactions can be engineered, therefore into inorganic materials and crystals that incorporate Sn, Ge, or other heavy and highly polarizable central atoms from across the periodic table. A careful selection of the substituents on the M center (along with the halide involved with the X-bond) can allow us to achieve interaction energies on par with or exceeding those observed for organic molecules. We hope to have opened up and started to answer fundamental questions about the nature of halogen bonds. The electrostatic (vs. covalent) nature of these forms of interactions continue to attract attention in the literature.3 The nature of the r-dependence of the interaction energies reveal for us the extent to which X-bonds deviate in nature from covalent interactions, simple dispersion interactions, and from each other based on the identity of Y. Universal rules for predicting the strength of halogen bonds for specific bases (aromatic or not, with one lone pair or two) may have to wait. But we have been gratified to find that functions of the simple form ΔE (or H) = A/rn + C link the geometry and the thermodynamics of halogen bonds formed by MF3I species with whole sectors of Lewis bases as distinct in composition as O(CH3)2 and NBr3. Supporting Information. Graphics of certain non-halogen bonded complexes, tables of interaction energy, enthalpy, and free energy data, and coordinates for the structures described in this work. This material is available free of charge via the Internet at http://pubs.acs.org. Acknowledgment Our work was supported by the National Science Foundation (NSF-CAREER award (CHE-1056430) and NSF-MRI Grants (CHE-0958696 (University of Richmond (UR)), and CHE-1229354 (the MERCURY consortium). 26


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References (1) See Odd Hassel’s 1969 Nobel lecture for examples of apparently halogen bonded complexes that were identified in that context as charge transfer complexes: “Structural Aspects of Interatomic Charge-Transfer Bonding”. Available free of cost at http://www.nobelprize. org/nobel_prizes/chemistry/laureates/1969/hassel-lecture.pdf Nobelprize.org. Nobel Media AB 2014. Retrieved July 30, 2014. The lecture is also available elsewhere (see reference 2). (2) Hassel, O. Structural Aspects of Interatomic Charge-Transfer Bonding Science 1970, 170, 497-502. (3) The discussion continues in the literature on the significance of the charge transfer versus the electrostatic nature of the halogen bond. See, for example, references 4-7 below. We do not contribute directly in this work to that conversation; the term “sub-covalent” affirms the presence of both influences and the fact that an A---B halogen bond is necessarily weaker than the corresponding pure A—B covalent bonds. (4) Legon, A. C. The Halogen Bond: an Interim Perspective. Phys. Chem. Chem. Phys. 2010, 12, 7736-7747. ( 5 ) Palusiak, M. On the Nature of Halogen Bond – The Kohn–Sham Molecular Orbital Approach. J. Mol. Struct. THEOCHEM, 2010, 945, 89-92. (6) Stone, A. J. Are Halogen Bonded Structures Electrostatically Driven? J. Am. Chem. Soc. 2013, 135, 7005-7009. (7) Wang, C.; Danovich, D.; Mo, Y.; Shaik, S. On The Nature of the Halogen Bond. J. Chem. Theory Comput. 2014, 10, 3726–3737. (8) Brinck, T.; Murray, J.S.; Politzer P. Surface Electrostatic Potentials of Halogenated Methanes as Indicators of Directional Intermolecular Interactions. Int. J. Quantum Chem. 1992, 44, 5764.

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(9) Murray, J.S.; Paulsen, K.; Politzer, P. Molecular Surface Electrostatic Potentials in the Analysis of Non-Hydrogen-Bonding Noncovalent Interactions. Proc. Indian Acad. Sci. (Chem. Sci.) 1994, 106, 267-275. ( 10 ) Clark, T.; Hennemann, M.; Murray, J.S.; Politzer, P. Halogen bonding: the σ-hole (Proceedings of “Modeling interactions in biomolecules II”, Prague, September 5th–9th, 2005). J. Mol. Model. 2007, 13, 291-296. (11) Emsley, J. The Elements, 3rd Ed., Oxford Univ. Press, Oxford, 1998. (12) Kutzelnigg, W. Chemical Bonding in Higher Main Group Elements. Angew. Chem. Int. Ed. Engl. 1984, 23, 272-295. (13) Politzer, P.; Harris, R.R. Properties of Atoms in Molecules. I. Proposed Definition of the Charge on an Atom in a Molecule J. Amer. Chem. Soc. 1970, 92, 6451-6454. (14) Politzer, P.; Lane, P.; Concha, M. C.; Ma, Y; Murray, J. S. An Overview of Halogen Bonding. J. Mol. Model 2007, 13, 305-311. (15) Politzer, P.; Murray, J. S. Halogen Bonding: an Interim Discussion. ChemPhysChem 2013, 14, 278-294. (16) Politzer et al. point out in ref. (15) and in references therein that halogen bonding is fundamentally a sub-group of the broader group of so-called σ-hole bonding interactions. His unifying argument harmonizes an entire range of interactions and opens up room too for the prediction possibly of other σ-hole based interactions that may have interesting applications. ( 17 ) Auffinger, P.; Hays, F.A.; Westhof, E.; Shing-Ho P. Halogen Bonds in Biological Molecules. Proc. Nat. Acad. Sci. 2004, 101, 16789-16794.

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(18) De Moliner, E.; Brown N. R.; Johnson, L. N. Alternative Binding Modes of an Inhibitor to two Different Kinases. Eur. J. Biochem. 2003, 270, 3174-3181. (19) Metrangolo, P.; Neukirch, H.; Pilati, T.; Resnati, G. Halogen Bonding Based Recognition Processes: A World Parallel to Hydrogen Bonding. Acc. Chem. Res. 2005, 38, 386-395. ( 20 ) Metrangolo, P.; Pilati, T.; Resnati, G.; Stevenazzi, A. Halogen Bonding Driven SelfAssembly of Fluorocarbons and Hydrocarbons Curr. Opin. Colloid Interface Sci. 2003, 8, 215-222. (21) Donald, K.; Wittmaack, B. K.; Crigger, C. Tuning Sigma-Holes: Charge Redistribution in the Heavy (Group 14) Analogues of Simple and Mixed Halomethanes can Impose Strong Propensities for Halogen Bonding. J. Phys. Chem. A 2010, 114, 7213-7222. ( 22 ) Head-Gordon, M.; Head-Gordon, T. Analytic MP2 Frequencies without Fifth-Order Storage. Theory and Application to Bifurcated Hydrogen Bonds in the Water Hexamer. Chem. Phys. Lett. 1994, 220, 122-128 and references therein. (23) Dunning Jr., T. H. Gaussian Basis Sets for use in Correlated Molecular Calculations. I. The Atoms Boron through Neon and Hydrogen J. Chem. Phys. 1989, 90, 1007-1023. (24) The correlation consistent triple-zeta (cc-pVTZpp) basis sets used are from the website of the Institute for Theoretical Chemistry at the University of Stuttgart: http://www.theochem. uni-stuttgart.de/pseudopotentials/clickpse.en.html. (25) Peterson, K. A. Systematically Convergent Basis Sets with Relativistic Pseudopotentials. I. Correlation Consistent Basis Sets for the Post-d Group 13–15 Elements J. Chem. Phys. 2003, 119, 11099-11112 (for Sn and Pb).

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( 26 ) Peterson, K. A.; Shepler, B. C.; Figgen, D.; Stoll, H. On the Spectroscopic and Thermochemical Properties of ClO, BrO, IO, and Their Anions J Phys. Chem. A, 2006, 110, 13877-13883 (for I). (27) Gaussian 09, Revision A.1, 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, Inc., Wallingford CT, 2009. (28) For a description of this procedure see Jensen, F. Introduction to Computational Chemistry; Wiley: New York, 1999; pp 172-173. ( 29 ) The Gaussview graphics program has been used to plot the representations of the electrostatic potentials and to visualize the molecular orbitals in this report. (30) The ChemCraft graphical program: http://www.chemcraftprog.com (last accessed August 1, 2014). (31) The actual values for the bond angles plotted in the figure are all included in the supporting information. (32) See the article “Defining a New Carbon Bond”, J. Kemsley, Chem and Eng. News, 2014, 92, 25-26. (33) Mani, D.; and Arunan, E. The X–CY (X = O/F, Y = O/S/F/Cl/Br/N/P) ‘Carbon Bond’ and Hydrophobic Interactions. Phys. Chem. Chem. Phys., 2013, 15, 14377-14383. (34) Bauá, A.; Mooibroek, T. J.; and Frontera, A. Tetrel-Bonding Interaction: Rediscovered Supramolecular Force? Angew. Chem. Int. Ed. 2013, 52, 12317-12321. ( 35 ) Thomas, S. P.; Pavan, M. S.; Guru Row, T. N. Experimental Evidence for ‘Carbon Bonding’ in the Solid State from Charge Density Analysis. Chem. Commun., 2014, 50, 4951. 30


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(36) Grabowski, S. J. Tetrel Bond–σ-Hole Bond as a Preliminary Stage of the SN2 Reaction. Phys. Chem. Chem. Phys., 2014, 16, 1824-1834. (37) Politzer, P.; Murray, J.; Clark, T. Halogen Bonding and Other σ-Hole Interactions: a Perspective. Phys. Chem. Chem. Phys., 2013, 15, 11178-11189. (38) Donald, K. J.; Tawfik, M. The Weak Helps the Strong: Sigma-Holes and the Stability of MF4Base Complexes J. Phys. Chem. A. 2013, 117, 14176-14183. (39) This result suggests that halogen bonds formed by R3Si—I are comparable to (that is, as weak as) some of those formed by R3C—Cl. This is hardly surprising since R3Si—I is known to form in general the weakest of the R3M—I---Y halogen bonds, where M is any group 14 atom, and a R3C—Cl---Y bond is inevitably weaker (because of the weaker Cl σhole) than the corresponding R3C—I---Y complex. (40) Wang, Z.-X.; Zhang, J.-C.; Wu, J.-Y.; Cao, W.-L. Theoretical Study on Intermolecular Interactions between Furan and Dihalogen Molecules XY ( X , Y = F , Cl , Br ) J. Chem. Phys. 2007, 126, 134301:1-7. (41) Cooke, S. A.; Corlett, G. K.; Holloway, J. H.; Legon, A. C. Evidence Concerning the Relative Nucleophilicities of Non-Bonding and -Bonding Electrons in Furan from the Rotational Spectrum of Furan···ClF J. Chem. Soc., Faraday Trans., 1998, 94, 2675-2680. See also related studies of interactions between ethene and terminal halide centers by the same research group (references 42, and 43). (42) Legon, A. C.; Thumwood, J. M. A. A π-Electron Donor–Acceptor Complex C2H4···Br2 Characterized By Its Rotational Spectrum Phys. Chem. Chem. Phys., 2001, 3, 1397-14202. (43) Stephens, S. L.; Mizukami, W.; Tew, D. P.; Walker, N. R.; Legon, A. C. The Halogen Bond between Ethene and a Simple Perfluoroiodoalkane: C2H4···ICF3 Identified by Broadband Rotational Spectroscopy J. Mol. Spec. 2012, 280, 47–53. 31



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(44) Wang, Z.; Zheng, B.; Yu, X.; Li, X.; Yi, P. Structure, Properties, And Nature Of The Pyridine - XY (X, Y = F, Cl, Br) Complexes: An Ab Initio Study J. Chem. Phys. 2010, 132, 164104:1-5. (45) Murray, J. S.; Lane, P.; Politzer, P. Expansion of the -hole Concept. J. Mol. Model. 2009, 15, 723-729. (46) Politzer, P.; Murray, J. S.; Lane, P.; Concha, M. C. Electrostatically Driven Complexes of SiF4 with Amines. Int. J. Quantum Chem. 2009, 109, 3773-3780. (47) Caminati, W.; Maris, A.; Dell'Erba, A.; Favero. P. G. Dynamical Behavior and DipoleDipole Interactions of Tetrafluoromethane-Water. Angew. Chem. Int. Ed. 2006, 45, 67116714. (48) See ref. 47 for an experimental investigation of the CF4H2O complex. A stabilizing dipoledipole type interaction between the C—F and O—H bonds (even though the two bonds are not necessarily parallel) is seen in that work to drive the formation of the complex. This is consistent with our observations in ref. 38 that the hydrogen bonding and bond-dipole interactions more generally dominate in the CF4H2O and CF4NH3 complexes, unlike in the SiF4Y complexes, for example, and the analogous complexes of the larger group 14 atoms where dative, sigma-hole supported, interactions are much more favorable. (49) Pauling, L. The Dependence of Bond Energy on Bond Length. J. Phys. Chem. 1954, 58, 662-666.

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