Molecular orbital theory explains chemical bonding by treating electrons as spread out over an entire molecule, not locked between just two atoms. When atomic orbitals overlap, they combine to form molecular orbitals with new energies and shapes. This model helps predict bond strength, magnetism, and electron arrangement in molecules.
It is especially important for molecules that simpler Lewis structures cannot fully explain, such as O2.
In an MO energy diagram, atomic orbitals from each atom are placed on the sides and the molecular orbitals formed from them are placed in the center. Constructive overlap creates lower energy bonding orbitals, while destructive overlap creates higher energy antibonding orbitals. Electrons fill molecular orbitals from lowest to highest energy following the Aufbau principle, Pauli exclusion principle, and Hund's rule.
The difference between bonding and antibonding electrons gives bond order, which predicts whether a bond is stable and how strong it is.
Key Facts
- Atomic orbitals combine to form molecular orbitals that extend over the whole molecule.
- Constructive overlap forms a bonding orbital with lower energy.
- Destructive overlap forms an antibonding orbital with higher energy and a node between nuclei.
- Bond order = (bonding electrons - antibonding electrons) / 2.
- A bond order greater than 0 predicts a stable bond, while bond order 0 predicts no stable bond.
- O2 is paramagnetic because it has two unpaired electrons in antibonding pi molecular orbitals.
Vocabulary
- Molecular orbital
- A molecular orbital is a region of space in a molecule where an electron is likely to be found.
- Bonding orbital
- A bonding orbital is a lower energy molecular orbital that increases electron density between nuclei and stabilizes a molecule.
- Antibonding orbital
- An antibonding orbital is a higher energy molecular orbital that has a node between nuclei and weakens bonding.
- Bond order
- Bond order is a measure of bond strength calculated from the number of bonding and antibonding electrons.
- Paramagnetism
- Paramagnetism is attraction to a magnetic field caused by one or more unpaired electrons.
Common Mistakes to Avoid
- Counting only valence electrons from one atom, which gives the wrong MO filling because the molecule contains electrons from both atoms.
- Putting paired electrons into equal energy orbitals too soon, which violates Hund's rule because degenerate orbitals fill singly before pairing.
- Treating antibonding electrons as neutral, which is wrong because they reduce bond order and weaken the molecule.
- Using the same 2p orbital order for every diatomic molecule, which can be wrong because B2, C2, and N2 have a different pi and sigma ordering than O2 and F2.
Practice Questions
- 1 H2 has 2 electrons in a bonding sigma 1s orbital and 0 electrons in an antibonding sigma 1s orbital. Calculate the bond order.
- 2 O2 has 10 bonding electrons and 6 antibonding electrons in its valence molecular orbitals. Calculate its bond order and state whether it should be stable.
- 3 Explain why molecular orbital theory predicts that O2 is paramagnetic, while a simple Lewis structure might suggest that all electrons are paired.