Types of bonding
Ionic bonding is electrostatic attraction between oppositely charged ions in a giant lattice; smaller, more highly charged ions attract more strongly. Metallic bonding is attraction between cations and delocalised electrons. A covalent bond is a shared pair attracted to both nuclei; in a dative (coordinate) bond both electrons come from one atom (, ). Head-on orbital overlap gives a σ bond; sideways p-orbital overlap gives a π bond. Period 3 atoms can expand the octet (, ).
≈ 45 minShapes of molecules and polarity
VSEPR theory: electron pairs around a central atom repel and get as far apart as possible; lone pairs repel more than bonding pairs, closing bond angles by about 2.5° each. 2 pairs: linear 180°; 3: trigonal planar 120°; 4: tetrahedral 109.5° (NH₃ pyramidal 107°, H₂O non-linear 104.5°); 5: trigonal bipyramidal 90°/120°; 6: octahedral 90°. Electronegativity is an atom's ability to attract the bonding electrons in a covalent bond; it increases across a period and up a group (F highest). Different electronegativities give polar bonds; a molecule is polar only if its dipoles don't cancel.
≈ 40 minIntermolecular forces
London (instantaneous dipole–induced dipole) forces exist between all molecules; they grow with the number of electrons and surface area. Permanent dipole–dipole forces act between polar molecules. Hydrogen bonds form when H is bonded to N, O or F and is attracted to a lone pair on N, O or F of another molecule; they are the strongest intermolecular force. Hydrogen bonding explains the high boiling point of water, NH₃ and HF, and why ice is less dense than water. Stronger intermolecular forces mean higher boiling points.
≈ 40 min