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 () ion, and the triiodide () 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. 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. " 2F−" 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