MO Theory and Covalent Bonding

Molecular orbital theory describes electron distribution throughout an entire molecule rather than confining electrons to specific bonds between atoms. This approach uses the linear combination of atomic orbitals to generate molecular orbitals where electrons occupy regions of space across the molecule. Unlike localized bonding models, molecular orbital theory accounts for delocalized electrons and provides better predictions of molecular stability and properties.

In-phase atomic orbital combinations create constructive interference, producing greater electron density between nuclei and forming lower-energy bonding molecular orbitals that stabilize molecules. Out-of-phase combinations create destructive interference with a nodal plane between nuclei, producing higher-energy antibonding molecular orbitals marked with an asterisk. The resulting electron density distribution directly determines molecular stability.

Sigma orbitals form from head-on overlap of atomic orbitals along the internuclear axis, such as s-s or end-to-end p-p overlap, with electron density centered on the axis. Pi orbitals result from sideways p-orbital overlap, concentrating electron density on opposite sides of the internuclear axis. Pi bonding and antibonding orbitals are typically degenerate, meaning they possess equal energy.

Bond order is calculated as the number of electrons in bonding orbitals minus electrons in antibonding orbitals, divided by two. A bond order greater than zero indicates that stable covalent bonds can exist between atoms. A bond order of zero means no stable bond forms, making this calculation essential for predicting molecular stability and bond strength.

The Lewis model cannot accurately represent benzene's delocalized electrons across the carbon ring. Molecular orbital theory assigns these electrons to three pi bonding molecular orbitals covering the entire hexagonal carbon structure, resulting in six electrons fully occupying bonding orbitals. This delocalization explains benzene's exceptional thermodynamic and chemical stability beyond what Lewis structures predict.

Higher electron density between positively charged nuclei increases electrostatic attraction and molecular stability. Bonding molecular orbitals concentrate electron density between nuclei, lowering energy and stabilizing molecules. Antibonding orbitals reduce electron density in the internuclear region, raising energy and destabilizing molecules. The net electron distribution determines overall molecular stability.

The extent and orientation of atomic orbital overlap determines the type and energy of resulting molecular orbitals. Head-on overlap along the internuclear axis produces sigma orbitals, while sideways overlap produces pi orbitals. Stronger overlap between atomic orbitals creates greater energy differences between bonding and antibonding orbitals, affecting bond strength and molecular properties.