Molecular Orbital Theory
Molecular Orbital (MO) Theory, developed by F. Hund and R.S. Mulliken in 1932, provides a sophisticated way of understanding molecular bonding by considering the wave nature of electrons. In this theory:
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Electrons in Molecules: Electrons in a molecule occupy molecular orbitals, which are analogous to atomic orbitals but are influenced by the nuclei of both atoms in the molecule.
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Combination of Atomic Orbitals: Atomic orbitals of similar energy and symmetry combine to form molecular orbitals. For each pair of atomic orbitals that combine, two molecular orbitals are produced: one bonding and one antibonding.
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Characteristics of Molecular Orbitals:
- Bonding Orbitals: These have lower energy than the atomic orbitals from which they are formed and stabilize the molecule.
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Antibonding Orbitals: These are higher in energy and destabilize the molecule due to the presence of a node between the nuclei.
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Energy Level Diagrams: Molecular orbitals are filled based on their energy levels, following the Pauli Exclusion Principle and Hund’s rule.
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Stability and Bond Order: The stability of the molecule can be assessed through its bond order, defined as half the difference between the number of electrons in bonding and antibonding orbitals. A higher bond order indicates more stability.
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Examples of Homonuclear Diatomic Molecules: The application of MO Theory is illustrated with examples such as
- Hydrogen (H2): Forms a bonding molecular orbital (C3_{1s}) and has a bond order of 1.
- Oxygen (O2): Exhibits paramagnetic behavior due to unpaired electrons in its molecular orbitals.
- Nitrogen (N2): Demonstrates a bond order of 3 and is extremely stable.
Understanding Molecular Orbital Theory is essential for explaining the magnetic and stability properties of different molecular species and establishes a foundation for further exploration of chemical bonding.