Valence Bond Theory and Covalent Bonds

A scientific theory explains natural laws and predicts behavior based on data.
Lewis theory visualizes electrons as dots; shared pairs form bonds, aiding in predicting valence electron distribution and atomic connectivity.
VSEPR theory predicts 3D molecular shapes but does not explain bonding.
Quantum mechanics describes electron locations:
- S orbital: spherical
- P orbital: dumbbell-shaped
Atomic orbitals alone cannot explain electron regions in molecules.

Valence bond theory defines a covalent bond as:
1. Overlap of half-filled atomic orbitals from two atoms
2. Formation of shared electron pairs from single electrons in overlapping orbitals
Bond strength depends on the overlap extent of orbitals; greater overlap = stronger bond.
Potential energy is linked to orbital overlap:
- Minimum potential energy occurs at bond distance.
- Energy to break bonds equals the energy decrease when bonds form.

Example: H–H bond breaking requires 7.24 × 10⁻¹⁹ J (or 4.36 × 10⁵ J per mole).
Stronger bonds have higher bond energies and lower potential energies; weaker bonds exhibit the opposite.
Average C–H bond energies:
- CH₄: 439.3 kJ/mol
- C₆H₅CH₃: 375.5 kJ/mol
- Average C–H: 413 kJ/mol