Valence Bond Theory

According to valence bond theory, a covalent bond forms when half-filled atomic orbitals from two atoms overlap, allowing their single electrons to combine into a shared electron pair. The mutual attraction between this negatively charged electron pair and the two positively charged nuclei creates the covalent bond. The strength of the bond depends on the extent of orbital overlap—greater overlap produces stronger bonds.

As hydrogen atoms move closer, their 1s orbitals begin to overlap. Initially, attractive forces between electrons and nuclei exceed repulsive forces, causing system energy to decrease. At an optimal distance of 74 picometers, energy reaches its minimum when attractive and repulsive forces balance. Further reduction in distance increases energy due to nucleus-nucleus and electron-electron repulsions, destabilizing the system.

Sigma (σ) bonds form from head-to-head orbital overlap along the internuclear axis, with electron density concentrated between nuclei. Pi (π) bonds result from side-by-side p orbital overlap, with electron density on opposite sides of the internuclear axis and a node along the axis itself. Single bonds are always σ bonds, while multiple bonds contain one σ bond and one or more π bonds.

Orbital orientation affects the degree of overlap between atomic orbitals. Orbitals oriented to overlap directly along the line between nuclei achieve greater overlap than misaligned orbitals. This direct alignment maximizes electron density concentration in the bonding region, producing stronger bonds. For example, end-to-end p orbital overlap creates stronger σ bonds than side-by-side arrangements.

Bond energy is the energy released when a bond forms, equal to the difference between separated atoms and the bonded state at optimal bond distance. Greater orbital overlap produces stronger bonds with higher bond energies. For instance, a carbon-carbon single bond averages 347 kJ/mol, while adding a π bond increases strength by 267 kJ/mol due to additional orbital overlap.

Two conditions are required: first, an orbital on one atom must overlap an orbital on a second atom, and second, the single electrons in each orbital must combine to form an electron pair. This shared electron pair, attracted to both positively charged nuclei, creates the covalent bond. The extent of orbital overlap determines bond strength and stability.

In multiple bonds, the first bond formed is always a σ bond, with remaining bonds being π bonds. A double bond consists of one σ and one π bond, while a triple bond contains one σ and two π bonds. Although π bonds increase overall bond strength, each individual π bond is generally weaker than a σ bond between the same atoms due to less extensive orbital overlap.