example, this orbital character does not appear for fluorine. As fluorine is very electronegative, it shares more of the sigma bonding electrons, creating a higher degree of sp hybridization than larger halogens. Moving additional electron density into the pz orbital affectively reduces the σ‐hole. For example, in a CF bond, 71.4% of electrons reside on F, whereas for less electronegative, larger halogens, like I, only ∼50% of the electron density resides on the halogen [8]. Meanwhile, the σ‐hole does not form for neutral, symmetric halogen containing molecules with equal electron distribution (e.g. carbon tetrahalides, hexahalobenzenes). This does not necessarily mean that symmetric or F‐based systems do not form halogen bonds; rather other attractive components become the dominate force.
1.4.5 Charge Transfer
Charge transfer has long been associated with halogen bonding, and Mulliken's investigations of I2 and organics containing O, S, or N heteroatoms are prime examples [42]. More recently, Palusiak utilized the Kohn–Sham molecular orbital (MO) theory to describe the interaction [148]. Halogen bonds and hydrogen bonds can have significant covalent character due to charge transfer from a guest to the antibonding σ* orbital (LUMO) of the R–X or R–H species [149] (Figure 1.15). The lower‐energy σ* orbital and higher‐energy σ orbital in this halogen bonding example allow for increased orbital mixing (σ orbital mixing shown for R–X donor in Figure 1.15b). These charge‐transfer adducts often result in lengthening of the RX or RH bond, which was highlighted in an early theoretical study of halogen bonding complexes between dihalogens (including interhalogens) and Lewis bases [150]. Here, elongation of the halogen–halogen bond is largest in the strongest complexes, up to 0.065 Å in the FBr⋯NH3 complexes. The study also demonstrated that the most polarizable halogens, and the interhalogens (FBr, FCl, etc.) with the biggest dipole, resulted in the largest interaction energies. Other studies have revealed that charge transfer can be a significant factor in organic halogen bond systems as well. One example evaluated complexes of bromocarbons (e.g. CBr3F, CBr3NO2, CBr3COCBr3, CBr3CONH2, Br3CCN) with anions (Br–, N3–, NCO–, and NCS–) [151,152]. In these reports, increasing charge transfer was linearly correlated with elongation of the CBr bond length. Therefore, as the interaction strength with the Lewis base increases, the CBr bond lengthens, suggesting that the donation of electrons to the antibonding σ* from the p‐orbital of the Lewis base results in a weakening of the CBr bond. These conclusions are further supported by MO theory where charge‐transfer effects are the leading component for organohalogen halogen bond formation in H3CX⋯O<span class="dbond"></span>CH2 and F3CX⋯O<span class="dbond"></span>CH2 (X = Cl, Br, I) models [148].
Figure 1.15 Simplified orbital‐interaction diagrams for (a) hydrogen‐bonded complexes DH⋯A− and (b) halogen‐bonded complexes DX⋯A− as they emerge from quantitative Kohn–Sham MO analyses.
Source: From Wolters and Bickelhaupt [149]. © 2012 John Wiley & Sons.
1.4.6 Dispersion and Polarization Component
London dispersion and polarization effects on the halogen bond can be important given the polarizability of larger halogens (e.g. I and Br) and the fact that interacting partners are frequently closer than the sum of the vdW radii. Dispersion interactions resulting from a temporary dipole and another dipole between the halogen and Lewis base provide a small but attractive force that contributes greatly in systems without large electropositive σ‐holes [153].
Lump–hole theory [154] is an alternative electrostatic model that describes a depletion of negative charge at the end of the halogen and accounts for dispersion and polarization. For example, Hobza and coworkers showed that a CH3Cl molecule can form a halogen bond with O<span class="dbond"></span>CH2 despite that the Cl never forms an electropositive σ‐hole [155]. Obviously, σ‐hole theory does not account for weak halogen bond formation in CH3Cl as dispersion dominates the interaction in this case.
Figure 1.16 DFT–SAPT decomposition analysis of H3CBr⋯NH3 (solid lines) and F3CBr⋯NH3 complexes (dashed lines) (kcal/mol). Potential energy minima are shown as vertical dashed lines (H3CBr⋯NH3 is green, and F3CBr⋯NH3 is light blue). The components are reported as follows: electrostatics (E(elec)), induction or polarization (E(ind)), dispersion (E(disp)), and exchange (E(exch)), and total binding energies (E(int)).
Source: From Riley and Hobza [153]. © 2013 Royal Society of Chemistry.
1.4.7 Decomposition
With decomposition analysis of intermolecular forces, contributions of electrostatics, induction or polarization, dispersion, and exchange repulsion are quantified. Decomposition of the halogen bond has allowed researchers to obtain a more complete view of the halogen bond. Symmetry‐adapted perturbation theory (SAPT) [156] and the density functional theory version (DFT–SAPT) [157] are used to describe the bonding components of the halogen bond. Total decomposition of the H3CBr⋯NH3 and F3CBr⋯NH3 halogen bond adducts (Figure 1.16) reveals notable differences between the two [153]. Specifically, the CH3 derivative was largely driven by inductive and dispersive forces, whereas the inclusion of CF3 groups led to a significantly larger electrostatic contribution.
1.4.8 Biological Computation of Halogen Bonding
Utilizing the halogen bond in biological settings is still novel. Currently, researchers are evaluating the influence of halogen bonding in protein stability, substrate binding, and drug design. Although nature seldom employs the halogen bond [158], medicinal chemists have found that the hydrophobicity of the halogen and directionality of the halogen bond could improve drug delivery and specificity. Drug design is time and cost intensive. To reduce this, medicinal chemists frequently turn to computational chemistry to identify target systems. However, specialized tools for modeling the halogen bond are still rare in the field.
Ho has been at the forefront of studying halogen bonding in biochemical systems. Using experiments, computations, and PDB searches, his group revealed that halogen bonds can stabilize ligand binding and molecular folding in proteins and nucleic acids [159]. An initial survey of the PDB found 113 different interactions when searching for short halogen‐Lewis base interactions. To date, more than 790 structures featuring the halogen bond in the PDB have been found [160]. A review by Ho et al. has also summarized current computational designs for halogen bonding drug candidates [161]. Using the structure of a protein and its binding pocket, their methodology identifies possible halogen bond acceptors within
