Molecular Orbital Theory — Chemistry Mnemonic
Target: mnemonic for molecular orbital theory bond order
Why is this hard to memorize?
Molecular Orbital Theory (MOT) explains bonding by combining atomic orbitals into molecular orbitals that belong to the entire molecule. Unlike VBT (hybridization), MOT can explain paramagnetism of O₂, bond orders of fractional values, and the existence/non-existence of molecules like He₂. For NEET and JEE, you need to know: the MO filling order (σ1s, σ*1s, σ2s, σ*2s, σ2p, π2p, π*2p, σ*2p), the bond order formula (BO = (Nb - Na)/2), and the MO diagrams for O₂, N₂, and simple diatomics.
Classic mnemonics you should know
"σ1s σ*1s σ2s σ*2s [π2px=π2py] σ2pz [π*2px=π*2py] σ*2pz"
For B₂, C₂, N₂ (up to 14 electrons): π2p comes BEFORE σ2p. For O₂, F₂ (15+ electrons): σ2p comes BEFORE π2p. The switch happens between N₂ and O₂ — this is a classic trick question.
"BO = (Bonding − Antibonding) / 2 = (Nb − Na) / 2"
Count electrons in bonding MOs (Nb) and antibonding MOs (Na). BO = (Nb−Na)/2. If BO > 0 → molecule exists. If BO = 0 → molecule doesn't exist (He₂). Higher BO → shorter, stronger bond.
"O₂ has 2 unpaired electrons in π*2p — that's why it's paramagnetic. VBT can't explain this!"
O₂ has 16 electrons. Filling MOs: the last 2 electrons go into the two degenerate π*2p orbitals, one each (Hund's rule). These 2 unpaired electrons make O₂ paramagnetic — attracted to magnets. This is MOT's greatest triumph.
The complete list
- Bond order = (Nb − Na) / 2
- σ bonding and σ* antibonding
- π bonding and π* antibonding
- MO order switches between N₂ and O₂
- O₂ is paramagnetic (2 unpaired e⁻)
- He₂ has BO = 0 (doesn't exist)
- N₂ has BO = 3 (triple bond)
Frequently asked questions
How do I calculate bond order using MOT?
BO = (Number of electrons in bonding MOs − Number in antibonding MOs) / 2. For H₂: 2 in σ1s(bonding), 0 in σ*1s(antibonding) → BO = (2−0)/2 = 1. For He₂: 2 in σ1s, 2 in σ*1s → BO = (2−2)/2 = 0 → He₂ doesn't exist as a stable molecule.
Why does the MO order change between N₂ and O₂?
In lighter elements (up to N), the 2s and 2p orbitals are close in energy, causing mixing (sp mixing) that pushes σ2p above π2p. In O₂ and beyond, the 2s-2p energy gap is larger, so mixing is negligible and σ2p drops below π2p (the "normal" order). This is why N₂ has σ after π, but O₂ has σ before π.
Why can VBT not explain O₂ paramagnetism?
Valence Bond Theory predicts O₂ as O=O with all electrons paired (diamagnetic). But experimentally, O₂ is attracted to magnets (paramagnetic) — meaning it has unpaired electrons. Only MOT correctly predicts this: the two electrons in the degenerate π*2p orbitals remain unpaired by Hund's rule.
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