2.1 Ionic Bonding & Born-Haber Cycle

Ionic bonds arise from electrostatic attraction between oppositely charged ions formed by electron transfer. Stability is captured by lattice enthalpy. The Born-Haber cycle applies Hess's law to compute lattice energy indirectly:

$U = \Delta H_f - \Delta H_{sub} - IE - \tfrac{1}{2}\Delta H_{diss} - EA$

NaCl example gives U = −787 kJ/mol. Higher lattice enthalpy → harder, higher-melting ionic solid.

2.2 Fajan's Rules

Fajan's rules quantify the partial covalent character of ionic bonds via cation-induced polarization of the anion:

• Small, highly charged cation → high polarizing power → more covalent character
• Large, easily polarizable anion → more covalent character
• Examples: $AlCl_{3}$ is more covalent than NaCl; CuCl is more covalent than KCl (same charge, but $Cu^{+}$ smaller and has 18-electron shell vs 8-electron shell for $K^{+}$ — more polarizing)

2.3 Covalent Bonding and Dipole Moment

Dipole moment μ = q × d (Debye). Polarity depends on both bond polarity and molecular geometry. Key: symmetric polyatomic molecules always have μ = 0 regardless of individual bond polarities. Non-symmetric molecules or those with lone pairs on the central atom have μ ≠ 0.

2.4 VSEPR Theory & Molecular Geometry

The central rule: electron pairs repel each other and adopt positions that minimize repulsion. Lone pairs occupy more space than bond pairs, compressing bond angles. The steric number (SN = σ bonds + lone pairs) determines the hybridization directly. Shape is determined after lone pairs are conceptually removed.

2.5 Hybridization

Hybridization is a mathematical model explaining the geometry of covalent compounds. Pure atomic orbitals are combined to form hybrid orbitals of equivalent energy and shape. The type of hybridization (sp → sp^{3}$$d^{2}) is determined entirely by SN. Hybridization determines geometry; geometry determines dipole moment.

2.6 Valence Bond Theory (Sigma & Pi Bonds)

VBT explains bond formation through orbital overlap. Sigma bonds (head-on) are present in every single, double, and triple bond. Pi bonds (lateral) accompany double and triple bonds. π bonds are weaker and prevent rotation, explaining cis-trans isomerism in alkenes.

2.7 Molecular Orbital Theory

MOT provides a more complete picture of bonding. Key results: bond order, magnetic character, and bond stability can all be derived from the MO electronic configuration. The critical test case is $O_{2}$, which MOT correctly identifies as paramagnetic. Bond order inversely relates to bond length and directly to bond energy.

2.8 Resonance, Hydrogen Bonding, Metallic Bonding

Resonance stabilizes molecules through electron delocalization. Hydrogen bonding (intermolecular in $H_{2}O$, HF; intramolecular in o-nitrophenol) explains anomalous properties. Metallic bonding (electron sea model) explains the electrical conductivity, malleability, and lustre of metals.