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CHEM 121 Studioby Learn4Less · UBC CHEM 121

CHEM 121 glossary

Valence bond theory

A model of bonding in which each covalent bond forms when an orbital on one atom overlaps an orbital on a neighbouring atom and the overlap region holds one shared pair of electrons with opposite spins. The electrons stay localized: in a bond between those two atoms, or in a lone pair on one atom. To explain molecular shapes, a central atom mixes its s and p orbitals into hybrid orbitals that point where VSEPR puts the electron domains; head-on overlap makes σ bonds and side-on overlap of the leftover p orbitals makes π bonds.

σ bond: head-on overlapTwo p orbitals on neighbouring atoms point straight at each other along the bond axis; they overlap in the region between the two nuclei, which is where the σ bond's electron density sits.bond axisABoverlap on the axiselectron density between the nuclei → σ bond(s + s, s + p and end-on p + p all give σ)

The VB picture of one bond: an orbital from each atom overlaps, and the overlap holds the shared pair.

Example

HX2\ce{H2}: each H brings one electron in a 1s1s orbital; the two 1s1s orbitals overlap and the shared pair is the H–H σ bond. CHX4\ce{CH4}: carbon mixes its 2s2s and three 2p2p orbitals into four sp3sp^3 hybrids at 109.5°, and each overlaps an H 1s1s orbital, giving four identical C–H σ bonds. The blind spot is OX2\ce{O2}: pairing every electron in bonds and lone pairs predicts a diamagnetic molecule, but OX2\ce{O2} is paramagnetic, which only MO theory explains.

See alsoHybridizationOrbital overlapσ bond (sigma bond)Molecular orbital theory