ChemistryChemical bondingClass 11
Molecular orbital theory: interactive simulation
Bond order is (bonding electrons − antibonding electrons) / 2. O₂ has two unpaired electrons in π*, so it is paramagnetic even though the Lewis structure looks like a complete double bond.
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Try this
- 1
Select O₂, then O₂⁺, then O₂⁻.
What you should see
Bond order goes 2 → 2.5 → 1.5. Removing a π* electron strengthens the bond; adding one weakens it.
- 2
Compare N₂ with O₂.
What you should see
N₂ has bond order 3 and is diamagnetic. O₂ has bond order 2 and two unpaired electrons.
Formulas behind it
Bond order
Common mistakes
- Using the O₂ orbital order for N₂. Up to N₂, π2p lies below σ2p.
FAQ
Why is O₂ paramagnetic?
The two π* orbitals each hold one electron, with parallel spins.
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