Ch 4 covers all types of chemical bonding — ionic, covalent, coordinate, hydrogen bonds — and theories for predicting molecular shape: VSEPR, valence bond (hybridisation), and molecular orbital theory.
Ionic bond: electron transfer (electronegativity difference > 1.7). Covalent: sharing (single, double, triple bonds). Coordinate: both electrons from one atom. Metallic: electron sea model. Lewis structures show bonding and lone pairs. Formal charge = valence e⁻ − lone pair e⁻ − ½ bonding e⁻. Resonance: multiple Lewis structures (O₃, CO₃²⁻).
VSEPR: molecular shape depends on electron domains (bonding + lone pairs). 2 = linear, 3 = trigonal planar, 4 = tetrahedral, 5 = trigonal bipyramidal, 6 = octahedral. Lone pairs distort angles (NH₃: 107°, H₂O: 104.5°). Hybridisation: sp (BeCl₂, linear, 180°), sp² (BF₃, trigonal, 120°), sp³ (CH₄, tetrahedral, 109.5°), sp³d, sp³d².
Atomic orbitals combine to form molecular orbitals: bonding (lower energy, σ, π) and antibonding (higher energy, σ*, π*). Bond order = ½(bonding − antibonding). O₂: bond order 2, paramagnetic (2 unpaired in π*). N₂: bond order 3, diamagnetic. Hydrogen bonding: strong intermolecular attraction (H-F, H-O, H-N), explains high bp of water.
Download: https://ncert.nic.in/textbook/pdf/kech104.pdf | Part I: https://ncert.nic.in/textbook/pdf/kech1ps.zip
In H₂O, oxygen is sp³ hybridised with 4 electron domains: 2 bonding pairs (O-H) and 2 lone pairs. VSEPR predicts a tetrahedral arrangement of domains, but the molecular shape (considering only atoms) is bent/angular at ~104.5°. Lone pairs occupy more space than bonding pairs, compressing the H-O-H angle below 109.5°.
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