Electron-deficient compound

An electron-deficient compound has too few valence electrons to describe all its bonds as ordinary two-center, two-electron bonds. It often uses multicenter bonding, in which electron pairs bind three or more atoms.

Three-center two-electron bond: A bond in which two electrons are shared across three atoms. This bonding motif lets electron-poor frameworks connect more atoms than ordinary bonds permit.
Valence electron: An electron in an atom's outer shell that can participate in chemical bonding. The available valence electrons determine whether conventional bonds can account for the framework.
Borane: A hydride of boron, many of which contain electron-deficient clusters or bridging hydrogen atoms. Boranes are the best-known family of compounds with electron-deficient bonding frameworks.
Electron-precise compound: A compound whose valence-electron count supports conventional localized two-center, two-electron bonds. It represents the ordinary bond-counting case that electron-deficient frameworks exceed.
Diborane: A boron hydride, B₂H₆, with two bridging hydrogen atoms and four terminal hydrogen atoms. Its B–H–B bridges are the classic example of three-center two-electron bonding.
Lewis structure: A diagram showing valence electrons as bonds and lone pairs around atoms. Lewis structures expose the electron shortfall when every connection is drawn as a conventional bond.
Carborane: A boron–carbon cluster compound whose cage structure is stabilized by delocalized skeletal bonding. Carboranes apply cluster bonding principles to cages containing both boron and carbon.
Electron-rich compound: A compound with more valence electrons than are needed for its conventional bonding framework. It contrasts a shortage of bonding electrons with excess electrons or nonbonding electron density.
Molecular orbital theory: A model of bonding that describes electrons as occupying orbitals spread across a molecule. Delocalized orbitals provide a direct account of bonding across three or more atoms.
Octet rule: A chemical rule of thumb that atoms tend to gain, lose, or share electrons to reach eight valence electrons. Electron-deficient compounds often depart from octet patterns without violating electron-counting principles.

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