Editing Hypervalent molecule (section)

==Bonding in hypervalent molecules==
Early considerations of the geometry of hypervalent molecules returned familiar arrangements that were well explained by the [[VSEPR model]] for atomic bonding.  Accordingly, AB<sub>5</sub> and AB<sub>6</sub> type molecules would possess a trigonal bi-pyramidal and octahedral geometry, respectively. However, in order to account for the observed bond angles, bond lengths and apparent violation of the Lewis [[octet rule]], several alternative models have been proposed.

In the 1950s an expanded valence shell treatment of hypervalent bonding was adduced to explain the molecular architecture, where the central atom of penta- and hexacoordinated molecules would utilize d AOs in addition to s and p AOs. However, advances in the study of ''[[ab initio]]'' calculations have revealed that the contribution of d-orbitals to hypervalent bonding is too small to describe the bonding properties, and this description is now regarded as much less important.<ref name="ReferenceA"/>  It was shown that in the case of hexacoordinated SF<sub>6</sub>, d-orbitals are not involved in S-F bond formation, but charge transfer between the sulfur and fluorine atoms and the apposite resonance structures were able to account for the hypervalency (See below).

Additional modifications to the octet rule have been attempted to involve ionic characteristics in hypervalent bonding. As one of these modifications, in 1951, the concept of the [[three-center four-electron bond|3-center 4-electron (3c-4e) bond]], which described hypervalent bonding with a qualitative [[molecular orbital]], was proposed. The 3c-4e bond is described as three molecular orbitals given by the combination of a p atomic orbital on the central atom and an atomic orbital from each of the two [[ligand]]s on opposite sides of the central atom. Only one of the two pairs of electrons is occupying a molecular orbital that involves bonding to the central atom, the second pair being non-bonding and occupying a molecular orbital composed of only atomic orbitals from the two ligands. This model in which the octet rule is preserved was also advocated by Musher.<ref name=Jensen/>  
[[image:XeF2.svg|thumb|400px|center|Qualitative model for a [[three-center four-electron bond]]]]

=== Molecular orbital theory ===
A complete description of hypervalent molecules arises from consideration of molecular orbital theory through quantum mechanical methods. An [[LCAO]] in, for example, sulfur hexafluoride, taking a basis set of the one sulfur 3s-orbital, the three sulfur 3p-orbitals, and six octahedral geometry symmetry-adapted linear combinations (SALCs) of fluorine orbitals, a total of ten molecular orbitals are obtained (four fully occupied bonding MOs of the lowest energy, two fully occupied intermediate energy non-bonding MOs and four vacant antibonding MOs with the highest energy) providing room for all 12 valence electrons.  This is a stable configuration only for S''X''<sub>6</sub> molecules containing electronegative ligand atoms like fluorine, which explains why SH<sub>6</sub> is not a stable molecule. In the bonding model, the two non-bonding MOs (1e<sub>g</sub>) are localized equally on all six fluorine atoms.

===Valence bond theory===
For hypervalent compounds in which the ligands are more [[electronegative]] than the central, hypervalent atom, [[resonance structures]] can be drawn with no more than four covalent electron pair bonds and completed with ionic bonds to obey the octet rule.  For example, in [[phosphorus pentafluoride]] (PF<sub>5</sub>), 5 resonance structures can be generated each with four covalent bonds and one ionic bond with greater weight in the structures placing ionic character in the axial bonds, thus satisfying the octet rule and explaining both the observed [[trigonal bipyramidal molecular geometry]] and the fact that the axial bond length (158 pm) is longer than the equatorial (154 pm).<ref>{{cite journal | title = A Simple Qualitative Molecular-Orbital/Valence-Bond Description of the Bonding in Main Group "Hypervalent" Molecules | journal = [[Journal of Chemical Education]] | year = 1998 | volume = 75 | pages = 910–915 | doi = 10.1021/ed075p910 | author1 = Curnow, Owen J. | issue = 7|bibcode = 1998JChEd..75..910C }}</ref>

[[image:penta phos.svg|thumb|500px | center | Phosphorus pentafluoride. There are 2 possible structures with an axial ionic bond, plus 3 possible structures with an equatorial ionic bond.]]

For a hexacoordinate molecule such as [[sulfur hexafluoride]], each of the six bonds is the same length.  The rationalization described above can be applied to generate 15 resonance structures each with four covalent bonds and two ionic bonds, such that the ionic character is distributed equally across each of the sulfur-fluorine bonds.

[[image:hexa sulf.svg|thumb|500px | center | Sulfur hexafluoride. There are 12 structures with the two ionic bonds in adjacent (''cis'') positions, plus 3 structures with the two ionic bonds in opposite (''trans'') positions.]]

Spin-coupled valence bond theory has been applied to [[diazomethane]] and the resulting orbital analysis was interpreted in terms of a chemical structure in which the central nitrogen has five covalent bonds;

[[File:Hypervalent diazomethane.png|thumb|150px| center|Chemical formula of diazomethane, showing hypervalent nitrogen]]

This led the authors to the interesting conclusion that "Contrary to what we were all taught as undergraduates, the nitrogen atom does indeed form five covalent linkages and the availability or otherwise of d-orbitals has nothing to do with this state of affairs."<ref>{{cite journal | title = Modern valence bond theory | journal = [[Chemical Society Reviews]] | year = 1997 | volume = 26 | pages = 87–100 | doi = 10.1039/CS9972600087 | author1 = Gerratt, Joe | issue = 2 }}</ref>