Molecular orbital theory
In molecular orbital (MO) theory, atomic orbitals combine to make orbitals spread over the whole molecule. Combining in phase gives a lower-energy bonding orbital. Out of phase gives a higher-energy antibonding orbital. Bond order = (bonding − antibonding electrons) / 2.
Last updated
Analogy
Two waves adding give a bigger wave (bonding). Two waves cancelling give a flat spot (antibonding).
Example
H₂: (2 − 0)/2 = 1He₂: (2 − 2)/2 = 0 (no bond)N₂: (10 − 4)/2 = 3O₂: (8 − 4)/2 = 2
In practice
MO theory explains why O₂ is attracted to a magnet, which a simple Lewis structure cannot.
Practice
Flashcards
- How is MO bond order calculated?(Bonding electrons − antibonding electrons) / 2.
Interview questions
- Use molecular orbital theory to explain why O₂ is paramagnetic and N₂ is not.
- O₂ has two electrons in separate π* orbitals
- These electrons are unpaired
- N₂ has no electrons in antibonding π*
- All N₂ electrons are paired
Questions people ask
Why does He₂ not exist as a stable molecule?
Two helium atoms bring four electrons. Two fill the bonding orbital and two fill the antibonding one. The antibonding electrons cancel the benefit of the bonding ones, so bond order is zero. The atoms gain nothing by joining, so they stay separate.
Why is O₂ paramagnetic?
In the MO diagram for O₂, the last two electrons go into two separate, equal-energy antibonding π orbitals, one each, with the same spin (Hund's rule). Two unpaired electrons make O₂ paramagnetic, so liquid oxygen sticks to a strong magnet.