Viewpoint articlePerspectives on halogen bonding and other σ-hole interactions: Lex parsimoniae (Occam’s Razor)
Graphical abstract
Molecular surface electrostatic potentials of SeCl2 reflect the anisotropies of the electronic densities of both the selenium (left) and chlorines (right) in this molecule; both atoms have regions of both positive (red, yellow, green) and negative (blue) electrostatic potential.
Highlights
► An overview of the “σ-hole” concept is given. ► σ-Holes on Groups IV–VII atoms explain many long known noncovalent interactions. ► Density difference plots of halogen bonding interactions are presented. ► Electrostatic potentials point to the fallacy of atomic charges. ► Electrostatics/polarization and dispersion can usually explain σ-hole bonding.
Section snippets
The σ-hole
It has long been known that many covalently-bonded atoms of Groups IV–VII can form noncovalent complexes with negative sites, e.g. the lone pairs of Lewis bases. For Group VII, this is called “halogen bonding,” although this term was first used (to our knowledge) only in 1978 [1]; such interactions of covalently-bonded halogen atoms were already reported long before that [2], [3], [4], [5], [6], [7], [8] and continue to be studied at a rapidly increasing rate in recent years. Analogous
The development and properties of σ-holes
What is the origin of a σ-hole? Consider first a neutral ground-state free atom. Its average electronic density is spherically-symmetrical, and the electrostatic potential created by its nucleus and electrons is positive everywhere; the contribution of the nucleus outweighs that of the dispersed electrons. This has been proven rigorously [39]. When the atom forms a covalent bond, this picture changes; rearrangement of electronic density takes place. Some will shift into the bonding region,
Noncovalent Interactions
The discovery that many covalently-bonded atoms of Groups IV–VII can have regions of both positive and negative electrostatic potential [24], [25], [26], [27], through which they can interact with negative and positive sites, respectively, explains a great many of the weak noncovalent complexes that these atoms have been observed to form [1], [2], [3], [4], [5], [6], [7], [8], [9], [10], [11], [12], [13], [14], [15], [16], [17], [18], [19], [20], [21], [22], [23]. If the interaction is with a
Decomposition of binding energies
There appears to be a widespread urge to decompose a noncovalent interaction into various “components” and to try to quantify their contributions to the binding energy. Among such supposed components are electrostatics, dispersion, charge transfer, polarization, exchange repulsion, induction, orbital interaction, Pauli repulsion, distortion, etc. Various procedures have been introduced for partitioning the binding energy into different subsets of this collection. (There is not universal
Electrostatics/polarization vs. charge transfer: much ado about nothing
Charge transfer from the negative site B to a σ* antibonding orbital of RA has sometimes been described as the driving force in the formation of a complex RA⋯B, or at least as occurring in conjunction with an electrostatic interaction. This merits further comment.
First, it needs to be emphasized that Mulliken’s focus in developing his charge transfer theory was upon the electronic transition from the ground state of a complex to an excited state, largely dative [68]; it was not upon the
A fallacy of atomic charges
It is very tempting to try to assign numerical charges to atoms in molecules, despite the absence of any rigorous basis for doing so. They are not physical observables and are not uniquely defined. A variety of procedures have been proposed: at least 30, already in 1994 [71]. There are correlations between some of these (although that is not necessarily a positive sign), but there can also be major disagreements; for instance, six different methods produced charges ranging from −0.478 to +0.564
σ-Hole bonding and hydrogen bonding
It has been established that halogen bonding in particular and σ-hole bonding in general can be competitive with hydrogen bonding [8], [30], [36], [77], [78]. It has in fact been suggested that the hydrogen bond is simply a special case of a σ-hole bond [34], [38], [51], [79]. The hydrogen σ-hole is again centered on the extension of the covalent bond to the hydrogen; however instead of the positive potential being relatively narrowly focused, as for typical σ-holes, it tends to be more
Summary
Our focus in this discussion has been upon the formation of “normal” σ-hole complexes with binding energies less than approximately 10–15 kcal/mol. It has been found that, depending upon the properties of the negative site and the magnitude of the σ-hole potential, much stronger interactions can occur [84], [85], having some degree of coordinate covalent character. While these are extremely interesting, they have not been addressed here.
Reed et al. report a quote from Gibbs [70]: “One of the
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