valence-bond theory
/ VAY-lens /
How does quantum mechanics describe a covalent bond without throwing away the comfortable picture of a bond as a line between two atoms? Valence-bond theory keeps that local picture: a covalent bond forms when an orbital on one atom, containing an unpaired electron, overlaps with an orbital on another atom, also containing an unpaired electron, and the two electrons pair up in the region of overlap. The more two orbitals overlap, the stronger the bond.
In this picture each bond is a localised pairing of electrons in overlapping atomic orbitals, sitting between just two atoms — exactly like the lines in a Lewis structure. To explain the observed shapes, valence-bond theory adds hybridisation: a central atom mixes its atomic orbitals (say one s and three p) into a set of equivalent hybrid orbitals (four sp3) that point toward the bonded atoms. Overlap along the line joining two nuclei gives a sigma bond; sideways overlap of leftover p orbitals gives pi bonds, which is how the theory accounts for double and triple bonds.
Valence-bond theory is the natural quantum-mechanical backbone for the everyday toolkit of Lewis structures, VSEPR and hybridisation, and it gives a wonderfully intuitive, localised account of molecular structure. Its honest limit is that, because it insists bonds are localised between two atoms, it struggles with delocalised electrons and with magnetic properties — most famously it predicts O2 to be diamagnetic, whereas O2 is in fact paramagnetic. For those cases molecular-orbital theory, with electrons spread over the whole molecule, is the better tool. The two theories are complementary, not rivals.
In methane, valence-bond theory says carbon's one 2s and three 2p orbitals mix into four equivalent sp3 hybrids, each overlapping with a hydrogen 1s orbital to make a sigma bond. Four equal bonds pointing to the corners of a tetrahedron fall straight out.
Overlap plus hybridisation: valence-bond theory builds methane's four equal bonds.
Valence-bond theory's localised bonds make it intuitive but limited: it famously mispredicts O2 as diamagnetic. For delocalisation and magnetism, molecular-orbital theory is needed; the two pictures complement each other.