Intermolecular forces become a properly ranked, universal set

Simple molecular substances have "weak intermolecular forces" between molecules, so they have low melting points. Often that is the entire treatment. Hydrogen bonding may be named for water.

Three named types, present in a defined hierarchy at comparable molar mass (Structure 2.2.9): London (dispersion) forces < dipole–dipole < hydrogen bonding. Crucially: London forces exist between all molecules and atoms, including non-polar ones, arising from instantaneous/induced dipoles, and they increase with molar mass and with electron cloud polarisability — which is why I₂ is a solid and F₂ a gas, and why LDFs in a large non-polar molecule can outweigh dipole–dipole forces in a small polar one. Hydrogen bonding requires H bonded to N, O or F and a lone pair on the acceptor. IB then uses IMFs to explain boiling points, solubility, viscosity, and chromatography (2.2.10) — separation by relative attraction to mobile vs stationary phase.

"Weak forces" becomes three distinguishable, rankable, explanatory entities. Also the intramolecular vs intermolecular distinction becomes an explicit exam trap: melting a molecular solid breaks IMFs, not covalent bonds.

GCSE needs only the contrast "giant vs simple molecular" to explain melting points.

The classic — "boiling water breaks the O–H bonds" (Taber and many others document this intramolecular/intermolecular conflation); assuming a polar molecule always boils higher than a non-polar one regardless of size; asserting hydrogen bonding in HCl or in CH₃OCH₃ (no O–H/N–H/F–H bond present); calling hydrogen bonding "a bond" of comparable strength to covalent.