Lesson 7 of 12 · 9 min
Valence bond theory and orbital overlap
NCERT § "Valence Bond Theory"
Lewis says two hydrogen atoms 'share a pair'. But electrons are spread-out clouds, not dots. What does sharing look like, and where does the 435.8 kJ/mol come from?
The lesson in notes
In short
Valence bond theory explains bond formation as the overlap of half-filled atomic orbitals holding electrons of opposite spin.
As two H atoms approach, attractions between each nucleus and the other's electron start to outweigh repulsions; energy falls to a minimum at 74 pm, where the H₂ molecule forms.
The greater the extent of overlap, the stronger the bond.
Overlap is positive (in-phase, same sign) when a bond forms, negative (out of phase) when it does not, and zero when orbital symmetry prevents interaction.
A sigma (σ) bond forms by end-to-end (head-on) overlap along the internuclear axis; it can involve s-s, s-p or p-p overlap.
A pi (π) bond forms by sideways overlap of parallel p orbitals, above and below the internuclear axis.
A σ bond is stronger than a π bond because head-on overlap is more extensive; a double bond is one σ plus one π, and a triple bond is one σ plus two π.
Simple overlap of pure orbitals cannot explain the equal bonds and 109.5° angles in CH₄, which led to the idea of hybridisation.