Topic 4 is often considered the most important IB Chemistry topic — bonding determines structure, which determines properties. Students learn to predict bonding type, draw Lewis structures, determine molecular shapes, and explain physical properties.
Ionic: electrostatic attraction between oppositely charged ions (metal + non-metal). Lattice structure: high melting point, conducts when molten/dissolved. Metallic: positive metal ions in a "sea" of delocalised electrons. Properties: electrical conductivity, malleability, high melting point.
Covalent: sharing of electron pairs between non-metals. Single (2e⁻), double (4e⁻), triple (6e⁻) bonds. Lewis structures: show all bonding and lone pairs. Steps: count valence electrons, place central atom, form bonds, distribute remaining electrons. Coordinate (dative) bonds: both electrons from one atom.
Electron domains (bonding + lone pairs) repel and arrange to maximise distance. Shapes: 2 domains → linear (180°), 3 → trigonal planar (120°), 4 → tetrahedral (109.5°), 5 → trigonal bipyramidal, 6 → octahedral. Lone pairs compress bond angles. Polar molecules: uneven charge distribution (shape + polar bonds).
London dispersion forces (all molecules, increase with electron count). Dipole-dipole (polar molecules). Hydrogen bonding (H bonded to F, O, or N — strongest IMF for small molecules). IMF strength determines boiling point, solubility, and viscosity. Giant covalent (diamond, SiO₂): very high melting points due to extensive covalent network.
Two conditions must both be met: (1) the molecule must contain polar bonds (difference in electronegativity between bonded atoms), AND (2) the molecular shape must be asymmetric so that bond dipoles do not cancel. CO₂ has polar C=O bonds but is linear → dipoles cancel → non-polar. H₂O has polar O-H bonds and is bent → dipoles do not cancel → polar.
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