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.
Understanding Chemistry: Molecular Orbital Theory
For two atomic orbitals to mix effectively, they need compatible shapes, the right orientation, and similar energies. Orbitals with very different energies usually mix only a little. This is why bonding in a molecule made from identical atoms differs from bonding in a molecule made from unlike atoms.
In hydrogen fluoride, fluorine attracts electrons more strongly than hydrogen. The molecular orbitals are therefore closer in energy and character to fluorine orbitals. This helps explain why the bond is polar and why electron density is not shared equally.
Orbital direction determines the type of bond formed. Head on overlap along the line joining two nuclei forms a sigma orbital. Sigma orbitals usually give strong bonds because the overlap is direct.
Sideways overlap of p orbitals forms pi orbitals. A double bond contains one sigma bond and one pi bond. A triple bond contains one sigma bond and two pi bonds.
Pi electrons sit above and below the bond axis, making them more exposed than sigma electrons. This affects chemical reactions. Many reactions occur at pi bonds because reagents can reach those electrons more easily.
Energy diagrams are not identical for every diatomic molecule. For molecules containing boron, carbon, or nitrogen, mixing between certain s and p orbitals changes the order of some molecular orbitals. For oxygen and fluorine, the ordering is different.
Students often memorize one diagram and apply it everywhere, which leads to wrong electron arrangements. First identify the atoms involved, then use the correct energy ordering. Count all valence electrons carefully, including charges for ions.
A positive charge means electrons have been removed. A negative charge means extra electrons have been added.
Molecular orbital theory connects directly to measurements in real chemistry. Bond order often matches observed bond length and bond energy. Higher bond order generally means a shorter, stronger bond, though real molecules can have exceptions.
The theory also explains magnetic behavior. A substance with unpaired electrons is attracted into a magnetic field, while one with only paired electrons is weakly repelled. This can be tested in a laboratory.
When reading an orbital diagram, pay close attention to nodes, electron arrows, and the energy level order. These details show whether electrons stabilize a molecule, weaken a bond, or produce magnetism.
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.