Molecular Orbital Theory (MOT)
Molecular Orbital Theory (MOT)
Molecular Orbital Theory (MOT) is a fundamental theory used to explain the electronic structure of molecules. Unlike Valence Bond Theory, which assumes that electrons are localized between atoms, MOT describes electrons in molecules as delocalized, where they occupy molecular orbitals that can extend over the entire molecule.
Key Concepts of MOT
- Molecular Orbitals (MOs): These are the orbitals that result from the combination of atomic orbitals (AOs) when atoms bond together. MOs are associated with the entire molecule rather than a single atom.
- Bonding and Antibonding Orbitals: When AOs combine, they form bonding MOs, which have lower energy and promote stability, and antibonding MOs, which have higher energy and can destabilize the molecule.
- Orbital Symmetry: For AOs to combine and form MOs, they must have compatible symmetry.
- Electron Configuration of Molecules: Electrons are filled into MOs according to the Aufbau principle, Hund's rule, and the Pauli exclusion principle.
Formation of Molecular Orbitals
When two atomic orbitals overlap, they can constructively interfere (in-phase) to form a bonding molecular orbital or destructively interfere (out-of-phase) to form an antibonding molecular orbital. The bonding MO is lower in energy than the original AOs, while the antibonding MO is higher in energy.
Energy Level Diagrams
Energy level diagrams are used to represent the relative energies of MOs. For diatomic molecules, the diagrams show the energy levels of the AOs of the separate atoms and how they combine to form the MOs of the molecule.
Molecular Orbital Diagrams for Homonuclear Diatomic Molecules
For homonuclear diatomic molecules (like H2, O2, N2), the MO diagrams are symmetrical. The s orbitals combine to form σ and σ* orbitals, while p orbitals combine to form π and π* orbitals as well as σ and σ* orbitals.
Molecular Orbital Diagrams for Heteronuclear Diatomic Molecules
For heteronuclear diatomic molecules (like CO, NO), the MO diagrams are not symmetrical due to the difference in electronegativity and energy levels of the AOs of the different atoms.
Bond Order
Bond order is an index of bond strength and is defined as the difference between the number of electrons in bonding MOs and antibonding MOs, divided by two.
[ \text{Bond Order} = \frac{\text{Number of electrons in bonding MOs} - \text{Number of electrons in antibonding MOs}}{2} ]
Paramagnetism and Diamagnetism
MOT can also explain the magnetic properties of molecules. Molecules with unpaired electrons in MOs are paramagnetic, while those with all electrons paired are diamagnetic.
Examples
Example 1: Molecular Orbital Diagram for O2
Oxygen has 8 electrons, and the MO diagram for O2 shows that there are two unpaired electrons in the π*2p orbitals. This explains the paramagnetic nature of O2.
Example 2: Bond Order Calculation for N2
Nitrogen has 7 electrons, so N2 has 14 electrons. The bond order for N2 can be calculated using the MO diagram and is found to be 3, indicating a triple bond.
Differences Between MOT and Valence Bond Theory
| Valence Bond Theory (VBT) | Molecular Orbital Theory (MOT) |
|---|---|
| Localized electron pairs | Delocalized electrons |
| Overlap of atomic orbitals | Combination of atomic orbitals into molecular orbitals |
| Does not explain paramagnetism of O2 | Explains paramagnetism and diamagnetism |
| No concept of antibonding orbitals | Includes antibonding orbitals |
| Less useful for polyatomic molecules | Applicable to both diatomic and polyatomic molecules |
Important Points
- MOT provides a more detailed and accurate description of the electronic structure of molecules than VBT.
- The theory can predict the magnetic properties and stability of molecules.
- MOT is particularly useful for understanding the bonding in complex molecules and ions.
By studying MOT, students can gain a deeper understanding of chemical bonding and the properties of molecules, which is essential for various applications in chemistry and materials science.