Molecular Orbital (MO) Theory explains the stability of molecules by describing how atomic orbitals combine to form molecular orbitals where electrons are shared across the entire molecule rather than just between two atoms.
Here’s a breakdown:
1. Formation of Molecular Orbitals
- When atoms come together, their atomic orbitals overlap and combine to form:
- Bonding orbitals (σ, π) → lower in energy than the original atomic orbitals.
- Antibonding orbitals (σ, π)** → higher in energy than the original atomic orbitals.
2. Electron Filling
- Electrons fill the molecular orbitals following:
- Aufbau principle (lowest energy first)
- Pauli exclusion principle (2 electrons per orbital, opposite spins)
- Hund’s rule (electrons fill degenerate orbitals singly first)
3. Bond Order and Stability
- Bond order = (electrons in bonding orbitals − electrons in antibonding orbitals) ÷ 2
- Positive bond order → stable molecule
- Bond order 0 → molecule does not exist
Example: Hydrogen molecule (H₂)
- 2 H atoms: each has 1 electron
- Combine 1s orbitals → σ(1s) bonding orbital and σ*(1s) antibonding orbital
- Both electrons go into σ(1s) → bond order = (2−0)/2 = 1 → H₂ is stable
4. Why It Explains Stability
- Electrons in bonding orbitals lower the overall energy of the molecule → stabilizes it.
- Electrons in antibonding orbitals increase energy → destabilize it.
- MO theory can also explain paramagnetism, diamagnetism, and the stability of molecules that Lewis structures cannot.
In short:
Molecular Orbital Theory explains stability by showing that more electrons in bonding orbitals than antibonding orbitals = a stable molecule, because bonding electrons lower the total energy.