Fill molecular orbital diagrams for homonuclear and heteronuclear diatomics.
Molecular orbital theory succeeds precisely where simple Lewis structures fail — most famously in explaining why O2 is paramagnetic despite textbook Lewis structures showing all electrons paired in a clean double bond. This makes MO theory a favorite AP Chemistry and IChO topic for testing genuine conceptual understanding over memorized rules.
When two atomic orbitals combine, they form two molecular orbitals: a lower-energy bonding orbital (constructive overlap, electron density concentrated between the nuclei) and a higher-energy antibonding orbital, marked with an asterisk (destructive overlap, a node between the nuclei). Sigma (σ) orbitals form from head-on overlap along the internuclear axis; pi (π) orbitals form from side-on overlap of parallel p orbitals, always occurring in pairs (πx and πy) at equal energy.
Bond order — a direct measure of bond strength and stability — is calculated as (bonding electrons - antibonding electrons)/2. For H2, both electrons occupy the bonding σ1s orbital, giving bond order 1 (a stable single bond); for He2, the bonding σ1s and antibonding σ1s* orbitals are both filled, giving bond order 0 — correctly predicting that He2 does not exist as a stable molecule, something Lewis theory has no mechanism to explain at all.
For period 2 diatomics, orbital energy ordering actually differs between the lighter elements (Li2 through N2) and the heavier ones (O2, F2, Ne2) due to s-p orbital mixing — a subtlety that changes which orbital fills first and is a common point of confusion. Filling the correct MO diagram for O2 reveals two unpaired electrons in the degenerate π2p* antibonding orbitals, correctly predicting paramagnetism (attraction to a magnetic field) that a simple Lewis structure entirely misses.
This MO diagram builder lets you fill molecular orbitals for any homonuclear diatomic from H2 through Ne2 (plus select heteronuclear examples), automatically calculates bond order, and flags paramagnetism from unpaired electrons — making visible exactly what Lewis structures cannot show.