In chemistry, a hypervalent molecule (the phenomenon is sometimes colloquially known as expanded octet) is a molecule that contains one or more main group elements apparently bearing more than eight electrons in their valence shells. Phosphorus pentachloride (), sulfur hexafluoride (), chlorine trifluoride (), the chlorite (ClO2–) ion in chlorous acid and the triiodide (I3–) ion are examples of hypervalent molecules.
Definitions and nomenclature Hypervalent molecules were first formally defined by Jeremy I. Musher in 1969 as molecules having central atoms of group 15–18 in any valence other than the lowest (i.e. 3, 2, 1, 0 for Groups 15, 16, 17, 18 respectively, based on the octet rule).
Several specific classes of hypervalent molecules exist: Hypervalent iodine compounds are useful reagents in organic chemistry (e.g. Dess–Martin periodinane) Tetra-, penta- and hexavalent phosphorus, silicon, and sulfur compounds (e.g. PCl5, PF5, SF6, sulfuranes and persulfuranes) Noble gas compounds (ex. xenon tetrafluoride, XeF4) Halogen polyfluorides (ex. chlorine pentafluoride, ClF5)
N-X-L notation N-X-L nomenclature, introduced collaboratively by the research groups of Martin, Arduengo, and Kochi in 1980, is often used to classify hypervalent compounds of main group elements, where: N represents the number of valence electrons X is the chemical symbol of the central atom L the number of ligands to the central atom Examples of N-X-L nomenclature include: XeF2, 10-Xe-2 PCl5, 10-P-5 SF6, 12-S-6 IF7, 14-I-7
History and controversy The debate over the nature and classification of hypervalent molecules goes back to Gilbert N. Lewis and Irving Langmuir and the debate over the nature of the chemical bond in the 1920s. | Link Lewis maintained the importance of the two-center two-electron (2c–2e) bond in describing hypervalence, thus using expanded octets to account for such molecules. Using the language of orbital hybridization, the bonds of molecules like PF5 and SF6 were said to be constructed from sp3dn orbitals on the central atom. Langmuir, on the other hand, upheld the dominance of the octet rule and preferred the use of ionic bonds to account for hypervalence without violating the rule (e.g. "SF42+ 2 F−" for SF6).
In the late 1920s and 1930s, Sugden argued for the existence of a two-center one-electron (2c–1e) bond and thus rationalized bonding in hypervalent molecules without the need for expanded octets or ionic bond character; this was poorly accepted at the time. In the 1940s and 1950s, Rundle and Pimentel popularized the idea of the three-center four-electron bond, which is essentially the same concept which Sugden attempted to advance decades earlier; the three-center four-electron bond can be alternatively viewed as consisting of two collinear two-center one-electron bonds, with the remaining two nonbonding electrons localized to the ligands.
The attempt to actually prepare hypervalent organic molecules began with Hermann Staudinger and Georg Wittig in the first half of the twentieth century, who sought to challenge the extant valence theory and successfully prepare nitrogen and phosphorus-centered hypervalent molecules. The theoretical basis for hypervalency was not delineated until J.I. Musher's work in 1969.
In 1990, Magnusson published a seminal work definitively excluding the significance of d-orbital hybridization in the bonding of hypervalent compounds of second-row elements. This had long been a point of contention and confusion in describing these molecules using molecular orbital theory. Part of the confusion here originates from the fact that one must include d-functions in the basis sets used to describe these compounds (or else unreasonably high energies and distorted geometries result), and the contribution of the d-function to the molecular wavefunction is large. These facts were historically interpreted to mean that d-orbitals must be involved in bonding. However, Magnusson concludes in his work that d-orbital involvement is not implicated in hypervalency.
Nevertheless, a 2013 study showed that although the Pimentel ionic model best accounts for the bonding of hypervalent species, the energetic contribution of an expanded octet structure is also not null. In this modern valence bond theory study of the bonding of xenon difluoride, it was found that ionic structures account for about 81% of the overall wavefunction, of which 70% arises from ionic structures employing only the p orbital on xenon while 11% arises from ionic structures employing an hybrid on xenon. The contribution of a formally hypervalent structure employing an orbital of sp3d hybridization on xenon accounts for 11% of the wavefunction, with a diradical contribution making up the remaining 8%. The 11% sp3d contribution results in a net stabilization of the molecule by mol−1, a minor but significant fraction of the total energy of the total bond energy ( mol−1). Other studies have similarly found minor but non-negligible energetic contributions from expanded octet structures in SF6 (17%) and XeF6 (14%).
Despite the lack of chemical realism, the IUPAC recommends the drawing of expanded octet structures for functional groups like sulfones and phosphoranes, in order to avoid the drawing of a large number of formal charges or partial single bonds.
Hypervalent hydrides A special type of hypervalent molecules is hypervalent hydrides. Most known hypervalent molecules contain substituents more electronegative than their central atoms. Hypervalent hydrides are of special interest because hydrogen is usually less electronegative than the central atom. A number of computational studies have been performed on chalcogen hydrides and pnictogen hydrides. Recently, a new computational study has shown that most hypervalent halogen hydrides XHn can exist. It is suggested that IH3 and IH5 are stable enough to be observable or, possibly, even isolable.
Criticism Both the term and concept of hypervalency still fall under criticism. In 1984, in response to this general controversy, Paul von Ragué Schleyer proposed the replacement of 'hypervalency' with use of the term hypercoordination because this term does not imply any mode of chemical bonding and the question could thus be avoided altogether.
The concept itself has been criticized by Ronald Gillespie who, based on an analysis of electron localization functions, wrote in 2002 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."
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 unknown.
The ionic model holds up well in thermochemical calculations. It predicts favorable exothermic formation of from phosphorus trifluoride PF3 and fluorine F2 whereas a similar reaction forming is not favorable.Predicting the Stability of Hypervalent Molecules Mitchell, Tracy A.; Finocchio, Debbie; Kua, Jeremy. J. Chem. Educ. 2007, 84, 629. Link
Alternative definition Durrant has proposed an alternative definition of hypervalency, based on the analysis of atomic charge maps obtained from atoms in molecules theory. This approach defines a parameter called the valence electron equivalent, γ, as "the formal shared electron count at a given atom, obtained by any combination of valid ionic and covalent resonance forms that reproduces the observed charge distribution". For any particular atom X, if the value of γ(X) is greater than 8, that atom is hypervalent. Using this alternative definition, many species such as PCl5, SO42–, and XeF4, that are hypervalent by Musher's definition, are reclassified as hypercoordinate but not hypervalent, due to strongly ionic bonding that draws electrons away from the central atom. On the other hand, some compounds that are normally written with ionic bonds in order to conform to the octet rule, such as ozone O3, nitrous oxide NNO, and trimethylamine N-oxide , are found to be genuinely hypervalent. Examples of γ calculations for phosphate PO43− (γ(P) = 2.6, non-hypervalent) and orthonitrate NO43− (γ(N) = 8.5, hypervalent) are shown below.
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, AB5 and AB6 type molecules would possess a trigonal bipyramidal 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 proposed, in which the central atom of penta- and hexacoordinated molecules was thought to utilize vacant d atomic orbitals in addition to its valence s and p orbitals to form hybrid orbitals. For example, phosphorus in was described as undergoing sp3d hybridization to accommodate five bonding pairs in a trigonal bipyramidal geometry, while sulfur in was treated as sp3d2 hybridized, consistent with an octahedral structure. This model provided a straightforward explanation within the valence bond framework for how atoms in the third period and beyond could exceed the octet rule by expanding their valence shells into the 3d subshell.
However, advances in ab initio quantum chemical calculations have suggested that the energetic contribution of d-orbitals to bonding in main group hypervalent molecules might be minimal. The high energy and poor radial overlap of the 3d orbitals with ligand orbitals result in negligible participation in bond formation. It was shown that in the case of hexacoordinated SF6, d-orbitals might not be significantly involved in S–F bond formation; rather, charge transfer between the central atom and ligands, along with appropriate resonance structures, can adequately explain the bonding characteristics and apparent hypervalency (see below). As a result, the d-orbital hybridization model is now regarded primarily as a historical or pedagogical tool.
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 (3c–4e) bond, which described hypervalent bonding using a qualitative molecular orbital framework, was proposed. The 3c–4e bond is described as three molecular orbitals formed by the combination of a p atomic orbital on the central atom with atomic orbitals from two ligands positioned linearly. Only one of the two pairs of electrons occupies a bonding orbital involving the central atom, while the second pair is nonbonding and delocalized between the two ligands. This model, which preserves the octet rule by distributing electrons across a delocalized system, was also later advocated by Musher.
Molecular orbital theory