🔗 Chemical Bonding and Molecular Structure

1. Kossel-Lewis Approach

Atoms achieve stable octet by transferring or sharing electrons. Octet Rule: Atoms tend to have 8 electrons in their valence shell.

Ionic Bond: Formed by the complete transfer of one or more electrons from one atom to another (electrostatic attraction). Favored by low ionization enthalpy of metal and high negative electron gain enthalpy of non-metal.
Covalent Bond: Formed by equal sharing of electrons between atoms.

2. VSEPR Theory

Valence Shell Electron Pair Repulsion Theory predicts the shape of molecules. Electron pairs (bond pairs and lone pairs) repel each other and try to stay as far apart as possible.

Order of Repulsion

Lone Pair - Lone Pair > Lone Pair - Bond Pair > Bond Pair - Bond Pair

Examples: BeCl₂ (Linear), BF₃ (Trigonal planar), CH₄ (Tetrahedral), NH₃ (Trigonal pyramidal due to 1 lone pair), H₂O (Bent due to 2 lone pairs).

3. Valence Bond Theory (VBT) and Hybridisation

A covalent bond is formed by the overlapping of half-filled atomic orbitals.

🧬 Hybridisation

The process of intermixing of atomic orbitals of slightly different energies to produce a set of entirely new equivalent orbitals (hybrid orbitals).

sp: Linear (180°) e.g., BeCl₂, Alkynes.
sp²: Trigonal planar (120°) e.g., BF₃, Alkenes.
sp³: Tetrahedral (109.5°) e.g., CH₄, Alkanes.

4. Molecular Orbital Theory (MOT)

Atomic orbitals combine to form molecular orbitals (MO). Addition of atomic orbitals forms Bonding MO (lower energy). Subtraction forms Antibonding MO (higher energy).

Bond Order = ½ (N_b - N_a) Where N_b = electrons in bonding MO, N_a = electrons in antibonding MO.

A positive bond order means a stable molecule. MOT successfully explains the paramagnetic nature of O₂.

5. Hydrogen Bonding

The attractive force which binds hydrogen atom of one molecule with an electronegative atom (F, O, or N) of another molecule.

Types: Intermolecular (between two molecules, e.g., H₂O) and Intramolecular (within the same molecule, e.g., o-nitrophenol).