Molecular orbital theory

Molecular orbital theory describes electrons as occupying orbitals spread across a molecule. It explains bonding through the formation and occupation of these orbitals from atomic orbitals.

Linear combination of atomic orbitals: A method for constructing molecular orbitals as weighted sums of atomic orbitals. It provides the mathematical recipe for combining atomic orbitals into molecular orbitals.
Diatomic nitrogen: A molecule consisting of two nitrogen atoms joined by a strong triple bond. Its filled bonding orbitals and empty antibonding orbitals account for its high bond order.
Atomic orbital: A quantum-mechanical state describing an electron's spatial distribution in an atom. Atomic orbitals are the starting functions used to construct molecular orbitals.
Valence bond theory: A bonding model that describes molecules using localized bonds formed by overlapping atomic orbitals. It localizes bonds where molecular orbital theory treats electrons as molecule-wide states.
Friedrich Hund: A German physicist who developed rules for electron configurations and molecular structure. His work helped establish how electrons occupy molecular energy levels.
Bonding orbital: A molecular orbital whose occupation lowers energy by concentrating electron density between nuclei. Constructive overlap creates the orbitals that stabilize molecules.
Oxygen: A chemical element whose common molecular form is a diatomic gas. Its two unpaired electrons in antibonding orbitals explain its paramagnetism.
Schrödinger equation: The equation governing quantum-mechanical wavefunctions and their energies. Molecular orbitals are approximate solutions to this equation for molecules.
Lewis structure: A diagram representing valence electrons as lone pairs and bonds between atoms. It offers a localized bookkeeping picture that molecular orbitals refine.
Robert S. Mulliken: An American physicist and chemist who developed molecular orbital theory and spectroscopy. He formalized and named the molecular orbital approach.

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