Chemical Bonding and Molecular Orbitals

Chemical bonding describes how shared electrons hold atoms together in molecules, and molecular-orbital theory explains it by combining atomic orbitals into delocalized molecular orbitals.

Definition

Chemical bonding is the net attractive interaction that holds atoms together in a molecule, arising quantum-mechanically from the lowering of electronic energy when electrons are shared; a molecular orbital is a one-electron wavefunction extending over the whole molecule, typically constructed as a linear combination of atomic orbitals.

Scope

This topic covers the quantum-mechanical description of the chemical bond: the hydrogen-molecule-ion and hydrogen-molecule as prototypes, the linear-combination-of-atomic-orbitals construction of bonding and antibonding molecular orbitals, bond order and its relation to stability, and the complementary valence-bond picture with hybridization. It explains covalent bonding, electron sharing, and the trends in bond strength and magnetism across diatomic molecules.

Core questions

• Why does sharing electrons lower the energy of two atoms and form a bond?
• How are molecular orbitals constructed from atomic orbitals?
• What distinguishes bonding from antibonding orbitals, and what is bond order?
• How do molecular-orbital and valence-bond descriptions relate?

Key concepts

• Covalent bond and electron sharing
• Bonding and antibonding orbitals
• Linear combination of atomic orbitals
• Bond order
• Exchange interaction
• Hybridization and resonance

Key theories

Molecular-orbital (LCAO) theory

Combining atomic orbitals in phase gives a bonding molecular orbital with enhanced electron density between the nuclei, and out of phase gives a higher-energy antibonding orbital with a node; filling these determines bond order and molecular stability.

Valence-bond theory and the covalent bond

Heitler and London showed quantum-mechanically that two hydrogen atoms bind through an exchange interaction of their electrons, the basis of the valence-bond picture later extended by Pauling with hybridization and resonance.

Clinical relevance

Molecular-orbital and valence-bond theories explain and predict the geometry, stability, reactivity, and magnetic and optical properties of molecules, providing the conceptual language of chemistry and the basis for computational methods used in drug discovery and materials design.

History

Heitler and London's 1927 treatment of the hydrogen molecule was the first quantum-mechanical explanation of the covalent bond. Mulliken and Hund developed the molecular-orbital approach in parallel, and Pauling's 1939 The Nature of the Chemical Bond synthesized valence-bond ideas with hybridization and electronegativity into a framework that transformed chemistry.

Key figures

• Walter Heitler
• Fritz London
• Robert Mulliken
• Linus Pauling

Seminal works

• heitler1927
• pauling1939
• atkins2011

Frequently asked questions

What is bond order?

Bond order is half the difference between the number of electrons in bonding and antibonding molecular orbitals. A higher bond order corresponds to a stronger, shorter bond; a bond order of zero means no stable bond forms, as for two helium atoms.

Why is molecular oxygen paramagnetic?

Molecular-orbital theory places the two highest electrons of O₂ in separate degenerate antibonding orbitals with parallel spins, leaving two unpaired electrons. This makes O₂ paramagnetic, a result the simple valence-bond picture does not naturally predict.