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A '''hypervalent molecule''' is a [[molecule]] that contains one or more typical elements ([[Periodic table group|group]] 1, 2, 13-18) formally bearing more than eight [[electron]]s in their [[valence shell]]s. [[Phosphorus pentachloride]] (PCl<sub>5</sub>), [[sulfur hexafluoride]] (SF<sub>6</sub>), the [[phosphate]] (PO<sub>4</sub><sup>3−</sup>) ion, [[Chlorine trifluoride]] and the [[triiodide]] (I<sub>3</sub><sup>−</sup>) 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.<ref>Musher, J.L. The Chemistry of Hypervalent Molecules ''Angew. Chem. Int. Ed. Engl.'' '''1969''', ''8'', 54-68. [http://dx.doi.org/10.1002/anie.196901311]</ref>

Several specific classes of hypervalent molecules exist:
* [[Hypervalent iodine]] compounds are useful reagents in organic chemistry.
* [[Sulfurane]]s and [[persulfurane]]s are hypervalent sulfur compounds.

==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. SF<sub>4</sub><sup>2+</sup>F<sub>2</sub><sup>2−</sup>). His opponent in this era, [[Gilbert N. Lewis]] on the other hand believed in '''octet expansion'''.<ref name=Jensen>''The Origin of the Term "Hypervalent" '' Jensen, William B. J. Chem. Educ. '''2006''' 83 1751. [http://jchemed.chem.wisc.edu/Journal/Issues/2006/Dec/abs1751.html Link]</ref>

Hypervalent bonding has also been believed to be described as sp<sup>3</sup>d and sp<sup>3</sup>d<sup>2</sup> [[Orbital hybridisation|hybrid orbital]]s 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{{Fact|date=January 2008}}.

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 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 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.<ref name=Jensen/>

==N-X-L notation==
The N-X-L nomenclature introduced in 1980.<ref>Perkins,C. W.; Martin, J. C.; Arduengo, A. J.; Lau, W.; Alegria, A,; Kochi, J. K.; An Electrically Neutral a-Sulfuranyl Radical from the Homolysis of a Perester with Neighboring Sulfenyl Sulfur: 9-S-3  species ''J.Am. Chem. Soc. 1980, 102, 7753-7759'' {{doi|10.1021/ja00546a019}}</ref> 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 
* [[xenon difluoride|XeF<sub>2</sub>]], '''10-Xe-2'''
* [[phosphorus pentachloride|PCl<sub>5</sub>]], '''10-P-5'''
* [[sulfur hexafluoride|SF<sub>6</sub>]], '''12-S-6'''
* [[iodine heptafluoride|IF<sub>7</sub>]], '''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.<ref name=Jensen/>

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."<ref>Gillespie, R. J.; Silvi, B. The octet rule and hypervalence: two misunderstood concepts. ''Coord. Chem. Rev.'' '''2002''', ''233-234'', 53-62. [http://dx.doi.org/10.1016/S0010-8545(02)00102-9]</ref>

For hypercoordinated molecules with [[electronegative]] ligands such as PF<sub>5</sub> 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 PF<sub>5</sub> do not have [[hydride]] counterparts e.g. [[phosphorane]] PH<sub>5</sub> which is an unstable molecule.

Even an ionic model holds up well in [[thermochemistry|thermochemical]] calculations. It predicts favorable [[exothermic]] formation of PF<sub>4</sub><sup>+</sup>F<sup>-</sup> from [[phosphorus trifluoride]] PF<sub>3</sub> and [[fluorine]] F<sub>2</sub> whereas a similar reaction forming PH<sub>4</sub><sup>+</sup>H<sup>-</sup> is not favorable.<ref>''Predicting the Stability of Hypervalent Molecules '' Mitchell, Tracy A.; Finocchio, Debbie; Kua, Jeremy. J. Chem. Educ. '''2007''', 84, 629. [http://jchemed.chem.wisc.edu/Journal/Issues/2007/Apr/abs629.html Link]</ref>

==References==
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[[Category:Chemical bonding]]

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