The physical properties of simple molecular substances are determined by the strength of their intermolecular forces, NOT their covalent bonds.
Boiling Points
To compare boiling points, apply this hierarchy:
- Hydrogen bonding present? (H–F, H–O, H–N) → anomalously high bp
- Polar molecule? → dipole-dipole forces increase bp vs non-polar of similar mass
- Molar mass / electron count? → more electrons = stronger London forces = higher bp
Solubility
Polar Solvents (H₂O)
Dissolve polar solutes & ionic compounds. Water forms ion-dipole or dipole-dipole interactions with the solute.
Non-polar Solvents (hexane)
Dissolve non-polar solutes via London forces. Cannot overcome strong ionic lattice or H-bonding networks.
Viscosity
Viscosity is the resistance to flow. Stronger / more extensive IMFs → higher viscosity. Long-chain hydrocarbons (more surface area → more London forces) are more viscous than short-chain ones.
⚠️ Examiner Trap – Boiling ≠ Breaking Bonds
When a substance boils, only intermolecular forces are overcome – the covalent bonds within each molecule remain intact. Never write "the covalent bonds break during boiling."
🔑 Bond Enthalpy – Gaseous State Only
Bond enthalpy values apply only to the gaseous state. If a substance is liquid or solid, you must first account for the energy needed to vaporise it. Data booklet values are averages across many compounds and may differ from actual values in specific molecules.