Chemistry
Molecular Orbital Theory for Diatomic Molecules
Quick fact
In molecular orbital theory, electrons are not assigned to individual bonds but are spread out over the whole molecule—this delocalization explains why O₂ is paramagnetic and why N₂ has an exceptionally strong triple bond.
Why this is interesting
You already know atoms stick together to form molecules, but have you ever wondered why oxygen gas is magnetic? The answer lies not in the atoms themselves, but in the invisible dance of electrons across the entire molecule.
Read the full explanation
Understanding Molecular Orbital Theory for Diatomic Molecules
Imagine two hydrogen atoms approaching each other. Each has a 1s orbital with one electron. As they get close, their atomic orbitals overlap and combine to form two new molecular orbitals: one lower in energy (bonding) and one higher in energy (antibonding). The two electrons fill the bonding orbital, lowering the system's energy and creating a stable H₂ molecule. For more complex diatomic molecules (like O₂, N₂, F₂), the process is similar but involves s and p orbitals. The resulting molecular orbitals can be sigma (end-to-end overlap) or pi (sideways overlap). Bond order—the net number of bonding pairs—is calculated as (bonding electrons - antibonding electrons)/2. A bond order of 1 means a single bond, 2 a double, and 3 a triple. The theory also predicts that if there are unpaired electrons in antibonding orbitals, the molecule will be attracted to a magnetic field (paramagnetic), as seen in O₂.
A deeper explanation
Molecular orbital theory reveals that electrons occupy orbitals that are quantum mechanical combinations of atomic orbitals, leading to constructive and destructive interference. The fundamental principle is the linear combination of atomic orbitals (LCAO), where the number of molecular orbitals equals the number of atomic orbitals combined. Bonding orbitals concentrate electron density between the nuclei, lowering energy, while antibonding orbitals have a node between the nuclei, raising energy. For diatomic molecules of second-row elements, the ordering of molecular orbital energies follows a specific pattern: σ1s, σ1s, σ2s, σ2s, then either σ2pz before the π2px,π2py pair (for B₂, C₂, N₂) or after (for O₂, F₂, Ne₂). This ordering depends on the energy difference between 2s and 2p orbitals. Understanding this theory is crucial because it explains why N₂ is inert (very high bond order of 3, no unpaired electrons) and why O₂ is reactive and paramagnetic (bond order 2, two unpaired electrons in π orbitals). It also correctly predicts that Be₂ and Ne₂ do not exist as stable diatomic molecules (bond order zero).