Hypervalent molecule

A hypervalent molecule is a molecule that contains one or more typical elements (group 1, 2, 13-18) formally bearing more than eight electrons in their valence shells. Phosphorus pentachloride (PCl₅), sulfur hexafluoride (SF₆), the phosphate (PO₄³⁻) ion, and the triiodide (I₃⁻) ion are examples of hypervalent molecules. Hypervalent molecules were first defined by Jeremy I. Musher in 1969 as molecules of the elements of group 15-18 in any oxidation state other than the lowest.^[1]

Several specific classes of hypervalent molecules exist:

• Hypervalent iodine compounds are useful reagents in organic chemistry.
• Sulfuranes and persulfuranes are hypervalent sulfur compounds.

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Bonding in hypervalent molecules

Since the hypervalent molecules do not seem to obey the octet rule, some models have been proposed to describe their bonding properties. Irving Langmuir in the 1920s held the view that the octet rule did prevail and that bonding was based on ionic interactions (e.g. SF₄²⁺F₂²⁻). His opponent in this era, Gilbert N. Lewis on the other hand believed in octet expansion.^[2]

Hypervalent bonding has also been believed to be described as sp³d and sp³d² hybrid orbitals composed of s, p, and d-orbitals at higher energy levels. 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 hybrid orbital description is now regarded as much less important.

As another description of hypervalent molecules, modifications of the octet rule have been attempted to involve ionic characteristics in hypervalent bonding. As one of these modifications, in 1951, the concept of the 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 orbital on a central atom and two ligand orbitals: an occupied bonding orbital, an occupied non-bonding orbital (HOMO), and an unoccupied anti-bonding orbital (LUMO). This model in which the octet rule is preserved was also advocated by Musher.^[2]

N-X-L notation

The N-X-L nomenclature introduced in 1980^[3]. is often used to classify hypervalent compounds of main group elements, where:

• N represents the number of valence electrons involved in bonding
• X is the chemical symbol of the central atom
• L the number of ligands to the central atom

Examples are

• XeF₂, 10-Xe-2
• PCl₅, 10-P-5
• SF₆, 12-S-6
• IF₇, 14-I-7

Criticism

As a replacement for the term hypervalency, Paul von Ragué Schleyer in 1984 proposed hypercoordination because this term does not imply any mode of chemical bonding.^[2]

The concept of hypervalency has been criticized by Ronald Gillespie who, based on an analysis of electron localization functions, concluded that "as there is no fundamental difference between the bonds in hypervalent and non-hypervalent (Lewis octet molecules) there is no reason to continue to use the term hypervalent."^[4]

For hypercoordinated molecules with electronegative ligands such as PF₅ it has been demonstrated that the ligands can pull away enough electron density from the central atom so that its net content is again 8 electrons or fewer. Consistent with this alternative view is the finding that hypercoordinated molecules based on fluorine ligands, for example PF₅ do not have hydride counterparts e.g. phosphorane PH₅ which is an unstable molecule.

Even an ionic model holds up well in thermochemical calculations. It predicts favorable exothermic formation of PF₄⁺F^- from phosphorus trifluoride PF₃ and fluorine F₂ whereas a similar reaction forming PH₄⁺H^- is not favorable.^[5]

References