Bonding Theory and Molecular Orbital Theory Notes

Bonding Theory

  • Requires quantum mechanics.
  • Two approximate methods of bonding:
    • Molecular Orbital (MO) Theory: Atomic orbitals are combined to form molecular orbitals.
    • Valence Bond (VB) Theory: Bonds are formed by the overlap of atomic orbitals.

Valence Bond Theory

  • Bonds are formed by overlap of atomic orbitals (AOs).
  • Example: H2 molecule
    • Lewis structure: H-H
    • Bond energy: 436 kJ/mol
    • Bond length: 74 pm (0.74 Å)
    • Formation: Overlap of 1s orbitals from each H atom.

Hybridization

  • Hybridization: Mixing of two or more atomic orbitals to create hybrid orbitals.
  • Number of hybrid orbitals = number of AOs combined.
  • Hybrid orbitals can have different shapes from the original AOs and are essential for accounting for molecular shapes.
  • Covalent bonds are formed by:
    • Overlap of hybrid orbitals with atomic orbitals.
    • Overlap of hybrid orbitals with other hybrid orbitals.

sp Hybrid Orbital

  • Example: BeH2, BeCl2
    • Molecular geometry: Linear
    • Hybridization involves mixing 2s and 2p orbitals to form sp hybrid orbitals.

sp² Hybrid Orbital

  • Example: BF3 molecule
    • Geometry: Trigonal planar
    • Hybridization of boron: sp² (33% s character, 67% p character).
    • Formation of 3 sp² orbitals by mixing 1 s and 2 p orbitals.

sp³ Hybrid Orbital

  • Example: CCl4
    • Geometry: Tetrahedral
    • Hybridization: sp³ (25% s character, 75% p character).
    • Formed by mixing 1 s and 3 p orbitals.

Multiple Bonds

  • Example: C2H4 (Ethylene)
    • Hybridization: sp²
    • Single bond (σ bond): Formed by head-on overlap.
    • p bond: Formed by sidewise overlap.

Bonding in Acetylene (C2H2)

  • Involves sp hybridization:
    • Ground state: 2s and 2p orbitals.
    • Hybridized orbitals contribute to a linear geometry.

Delocalized Electrons

  • Benzene (C6H6): Example of delocalized electrons.
    • Structure: p orbitals overlap sidewise around the ring.
    • No localized double bonds.

Molecular Orbital Theory

  • Molecular orbitals (MOs) are formed from the interaction of atomic orbitals.
  • The number of MOs formed equals the number of AOs combined.
  • Types of MOs:
    • Bonding orbitals: Lower energy than AOs, electron density overlaps.
    • Antibonding orbitals: Higher energy than AOs, electron density does not overlap.
  • Pauli’s Exclusion Principle: Each MO can hold two electrons.
  • Filling of MOs follows:
    • Low to high energy (Hund's rule applies for orbitals of the same energy).

Examples of Molecular Orbital Configurations

  • O2: Molecular orbital diagram shows it to be paramagnetic due to unpaired electrons.
  • N2: Bond order = 3 (indicating a triple bond).
  • General Principles:
    1. The MO number equals the AO number.
    2. Stable bonding MOs are less stable than corresponding antibonding MOs.
    3. Each MO can accommodate up to two electrons.
    4. Bond order calculation: \text{Bond order} = \frac{1}{2} (# \text{ of bonding electrons} - # \text{ of antibonding electrons}).

Magnetic Properties

  • Paramagnetic: Molecules with unpaired electrons (e.g., O2). Align in magnetic field.
  • Diamagnetic: Molecules with no unpaired electrons (e.g., N2). Do not align in magnetic field.

Heteronuclear Diatomic Molecules

  • Example: NO
    • Bond order calculation applies similarly, taking into account differing atomic energy levels.