Molecular Orbital Theory Overview

Overview

This lecture introduces Molecular Orbital (MO) Theory as a way to describe covalent bonding, compares it to Valence Bond Theory, and demonstrates how MO diagrams explain bond order, magnetism, and resonance.

Molecular Orbital Theory Concepts

• Valence Bond Theory describes bonding as the overlap of atomic orbitals.
• MO Theory treats atomic orbitals as combining to form new, communal molecular orbitals.
• Atomic orbitals are regions around an atom where electrons are likely found; molecular orbitals describe these regions for entire molecules.
• Two atomic orbitals form two molecular orbitals: one lower-energy bonding orbital (sigma) and one higher-energy antibonding orbital (sigma*).

Bonding and Antibonding Orbitals

• Bonding orbitals (sigma) have high electron probability between nuclei, stabilizing the molecule.
• Antibonding orbitals (sigma*) have a node (zero electron probability) between nuclei and destabilize the molecule.
• Electrons fill the lowest available energy orbitals first, favoring bonding orbitals.

Bond Order and Stability

• Bond order = (electrons in bonding orbitals – electrons in antibonding orbitals) ÷ 2.
• A bond order > 0 indicates a stable molecule; bond order = 0 indicates no bond forms.
• Bond order predicts number of bonds: single (1), double (2), triple (3), etc.

Examples

• H₂ (hydrogen): Bond order = 1 (stable single bond).
• He₂ (helium): Bond order = 0 (no bond forms, molecule unstable).
• O₂ (oxygen): Bond order = 2 (double bond with unpaired electrons, paramagnetic).
• N₂ (nitrogen): Bond order = 3 (triple bond, diamagnetic).
• NO (nitric oxide): Bond order = 2.5 (paramagnetic).

Molecular Orbital Diagrams for p Orbitals

• p orbitals combine to form sigma, pi, and their respective antibonding orbitals.
• The sequence of orbital energies can differ between molecules (e.g., in O₂, sigma bonding lower than pi bonding; in N₂, pi is lower than sigma).
• Electrons fill these orbitals based on increasing energy.

Magnetism and MO Theory

• MO Theory explains paramagnetism (unpaired electrons, attracted to magnetic fields) and diamagnetism (paired electrons, weakly repelled).
• O₂ is paramagnetic, which cannot be predicted by Valence Bond Theory.

Resonance and MO Theory

• MO Theory allows for delocalized electrons, explaining resonance structures better than Valence Bond Theory.

Key Terms & Definitions

• Atomic Orbital — A region around an atom where an electron is likely found.
• Molecular Orbital (MO) — A region in a molecule where an electron is likely found.
• Bonding Orbital (σ) — A molecular orbital with lower energy and increased electron density between nuclei.
• Antibonding Orbital (σ*) — A molecular orbital with higher energy and a node between nuclei.
• Bond Order — (Number of electrons in bonding MOs – number in antibonding MOs) ÷ 2.
• Paramagnetic — Species with unpaired electrons, attracted by magnetic fields.
• Diamagnetic — Species with only paired electrons, weakly repelled by magnetic fields.
• Resonance — Delocalization of electrons over more than two atoms, explained by MO Theory.

Action Items / Next Steps

• Practice drawing molecular orbital diagrams for diatomic molecules.
• Calculate bond order for given MO diagrams.
• Determine if molecules are paramagnetic or diamagnetic using MO configurations.
• Review textbook figures for nitrogen and oxygen MO diagrams.