Chemical Bonding and the Structure of Matter
Core concepts
Candidates must distinguish ionic, covalent (polar and nonpolar), metallic and intermolecular bonding, and connect each to physical properties. For covalent substances, use the octet rule, Lewis structures, VSEPR geometry, hybridisation and polarity of the whole molecule (not just the bonds). For ionic solids, know lattice energy and why melting points scale with charge and small ionic radius. For metals, the electron-sea model explains conductivity and malleability. Intermolecular forces - hydrogen bonds, dipole–dipole, London dispersion - explain boiling points, solubility and the anomalous behaviour of water and hydrogen fluoride.
Traps that appear in questions
- Bond polarity vs molecular polarity: CO₂ has polar bonds but is nonpolar because the linear geometry cancels the dipoles; water is bent and therefore polar.
- Hydrogen bonding: it requires H bonded to N, O or F. A molecule with H bonded to carbon does not hydrogen-bond as a donor.
- Melting-point comparisons: ionic solids usually melt higher than covalent molecular solids but lower than giant covalent networks such as diamond or silicon dioxide.
- Conductivity: ionic compounds conduct only when molten or dissolved; graphite conducts because of delocalised electrons; diamond does not.
- VSEPR shapes: lone pairs compress bond angles, so ammonia is about 107° and water about 104.5°, both less than methane's 109.5°.
Worked example
Compare the boiling points of HF, HCl and HBr. HCl and HBr rely on dipole–dipole and dispersion forces, which grow with electron count, so HBr > HCl. HF, despite the smallest mass, boils highest because of strong hydrogen bonding.
Rule of thumb
Draw the Lewis structure, apply VSEPR, then check whether the bond dipoles cancel by symmetry.