Molecular Orbital Theory is a concept in chemical bonding, and explain the behavior of electrons in molecules more effectively than earlier theories like the Valence Bond Theory. Molecular Orbital Theory describes the formation of molecular orbitals by the linear combination of atomic orbitals. These molecular orbitals extend over the entire molecule, and electrons in these orbitals are delocalised fairly than being confined between individual atoms. This theory is important for explaining phenomena such as bond order, magnetism, and bond energy in molecules.
Main Concepts of Molecular Orbital Theory
1. Atomic Orbitals and Molecular Orbitals: According to Molecular Orbital Theory, atomic orbitals of atoms involved in bonding combine to form molecular orbitals. When two atomic orbitals overlap, then two types of molecular orbitals formed:
Bonding Molecular Orbital: This orbital is formed by the constructive interference of atomic orbitals, where the electron density between the two nuclei increases, stabilising the molecule. The energy of the bonding orbital is lower than the energy of the atomic orbitals that combined.
Antibonding Molecular Orbital: This orbital is formed by destructive interference, where the electron density between the nuclei is reduced. Antibonding orbitals are higher in energy than the parent atomic orbitals and destabilise the molecule
2. Linear Combination of Atomic Orbitals (LCAO):
Molecular orbitals are formed by the linear combination of atomic orbitals. For this combination to occur, atomic orbitals must have similar energy levels and proper orientation. Only atomic orbitals are of similar energy and symmetry can combine to form molecular orbitals. For example, the 1s orbital of one atom can combine with the 1s orbital of another atom, but not with a 2p orbital due to their energy difference.
3. Molecular Orbital Diagrams: Molecular orbital diagrams are used to visualise the energy levels of the molecular orbitals formed from atomic orbitals. In a classic molecular orbital diagram:
Bonding orbitals are placed lower on the energy scale.
Antibonding orbitals are placed higher.
The atomic orbitals are shown on either side, with lines representing the combination into molecular orbitals in the middle. For diatomic molecules like hydrogen (H₂), oxygen (O₂), and nitrogen (N₂), molecular orbital diagrams explain the bonding and antibonding electrons.
4. Bond Order: Molecular Orbital Theory allows the calculation of bond order, which provides insight into the strength and stability of a bond. The bond order is calculated using the formula:
Bond Order = (nb−na) / 2
Where:
nb is the number of electrons in bonding molecular orbitals.
na is the number of electrons in antibonding molecular orbitals.
A positive bond order indicates a stable molecule, a bond order of zero suggests that the molecule is unstable or does not exist. For example, in H₂, there are two electrons in a bonding orbital and none in an antibonding orbital, giving a bond order of 1, indicating a stable single bond.
5. Magnetic Properties: Molecular Orbital Theory also explains the magnetic properties of molecules. If a molecule has unpaired electrons in its molecular orbitals, it exhibits paramagnetism, meaning it is attracted to a magnetic field. If all the electrons are paired, the molecule is diamagnetic and repelled by a magnetic field. For instance, oxygen (O₂) is paramagnetic because it has unpaired electrons in its antibonding orbitals, which cannot be explained by earlier theories like VBT.
6. Molecular Orbital Configurations for Simple Molecules:
Hydrogen Molecule (H₂): In H₂, two hydrogen atoms come together, and their 1s atomic orbitals combine to form a bonding molecular orbital (σ1s) and an antibonding molecular orbital (σ*1s). The two electrons of H₂ occupy the σ1s bonding orbital, resulting in a bond order of 1, which corresponds to a stable single bond.
Oxygen Molecule (O₂): In O₂, the atomic orbitals of the two oxygen atoms combine to form bonding and antibonding orbitals. The bond order in O₂ is 2, indicating a double bond. Moreover, O₂ has two unpaired electrons in its antibonding orbitals, making it paramagnetic.
Limitations of Molecular Orbital Theory
While Molecular Orbital Theory provides a more accurate description of molecular bonding, it has certain limitations:
It requires complex calculations for molecules with more than two atoms.
It does not provide a clear picture of localized bonds like sigma (σ) and pi (π) bonds, as described by Valence Bond Theory.
Applications of Molecular Orbital Theory
Molecular Orbital Theory is widely used in understanding and predicting the behavior of molecules, especially for:
Predicting molecular stability and reactivity.
Explaining the magnetic properties of molecules.
Determining bond order and bond energy.
Electronic structure of molecules in spectroscopy and quantum chemistry
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