Class 11 Chapter 3 Molecular Orbital Theory with Examples
Introduction
Molecular Orbital (MO) Theory describes covalent bonding by combining atomic orbitals to form molecular orbitals that extend over the entire molecule. Unlike Valence Bond Theory, MO Theory accounts for delocalization of electrons and explains properties such as bond order and magnetism.
Fundamentals of MO Theory
- Atomic orbitals combine in-phase to form bonding MOs and out-of-phase to form antibonding MOs.
- The number of MOs equals the number of original atomic orbitals.
- Electrons fill MOs according to Aufbau principle, Hund’s rule, and Pauli exclusion principle.
Bonding & Antibonding MOs
Bonding orbitals (σ, π) lower the energy and stabilize the molecule. Antibonding orbitals (σ*, π*) raise the energy and weaken the bond if occupied.
Molecular Orbital Energy Diagram
For diatomics up to N₂, the ordering is: σ(2s), σ*(2s), σ(2pz), π(2px,y), π*(2px,y), σ*(2pz). For O₂ and beyond, the σ(2pz) lies above π(2px,y).
Key Examples
H₂ Molecule
Two 1s orbitals combine to form σ(1s) and σ*(1s). Filling σ(1s) with two electrons gives bond order = (2 – 0)/2 = 1. H₂ is diamagnetic.
O₂ Molecule
Configuration: (σ2s)²(σ*2s)²(σ2pz)²(π2px,y)⁴(π*2px,y)². Bond order = (8 – 4)/2 = 2. O₂ is paramagnetic due to two unpaired electrons in π* orbitals.
CO Molecule
CO has 10 valence electrons. MO filling leads to bond order = 3. Despite unequal electronegativities, MO Theory predicts a small dipole moment and strong triple bond character.
Bond Order & Magnetic Properties
Bond order = (number of electrons in bonding MOs – number in antibonding MOs)/2. Higher bond order indicates stronger bond. Paramagnetism arises when unpaired electrons occupy antibonding MOs.
Conclusion
MO Theory provides a comprehensive framework for understanding molecular structure, bond strength, and magnetic behavior. Its delocalized approach explains phenomena that Valence Bond Theory cannot, making it essential for advanced chemical bonding studies.
