The big idea: Inside a molecule, atoms are held by strong covalent bonds. But molecules are also pulled towards each other by much weaker forces called intermolecular forces (IMFs).
It is these IMFs — not the covalent bonds — that you must overcome to melt or boil a molecular substance. So IMFs set the boiling point: stronger IMFs → more energy needed → higher boiling point.
A hydrogen bond (dashed) is the attraction between a δ+ hydrogen (bonded to O) of one water molecule and a δ− oxygen lone pair of another.
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Two different forces — don't mix them up: - Covalent bond (intramolecular) — holds atoms together inside a molecule. Strong. - Intermolecular force (between molecules) — pulls separate molecules towards each other. Much weaker.
Boiling water separates the H2O molecules from each other — it does not break the O–H covalent bonds.
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There are three types of intermolecular force. The order of increasing strength is the one you must know:
London (dispersion) < dipole–dipole < hydrogen bonding.
| Intermolecular force | Acts between | Relative strength |
|---|---|---|
| London (dispersion) | all molecules (only force in non-polar ones) | weakest |
| Dipole–dipole | polar molecules (permanent δ+/δ−) | stronger |
| Hydrogen bonding | molecules with H–N, H–O or H–F | strongest |
London (dispersion) forces: Present between all molecules, and the only force between non-polar molecules.
They arise from temporary, instantaneous dipoles: electrons move at random, so at any instant a molecule is slightly uneven, inducing a dipole in its neighbour.
More electrons → larger, more polarisable molecule → stronger London forces. This is why they get stronger down a group and as molecules get bigger.
Dipole–dipole forces: Act between polar molecules, which have a permanent dipole (a δ+ end and a δ− end from an electronegativity difference).
The δ+ end of one molecule is attracted to the δ− end of the next. For molecules of similar size, a polar molecule boils higher than a non-polar one.
Hydrogen bonding — the strongest: A special, extra-strong dipole–dipole force. It needs hydrogen bonded directly to N, O or F (the three most electronegative atoms).
The very δ+ hydrogen is strongly attracted to a lone pair on the N, O or F of a neighbouring molecule.
Memory hook: hydrogen bonding only happens with N, O, F — 'H bonds to NOF'.
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Almost every exam question here is the same skill: 'explain why X has a higher/lower boiling point than Y' — by comparing their intermolecular forces. There are two classic patterns.
Pattern 1 — boiling point rises down a group / with molecular size: Compare the noble gases or the halogens: boiling point rises as you go down.
Larger molecules have more electrons, so their London (dispersion) forces are stronger → more energy is needed to separate them → higher boiling point. The same logic explains why bp rises along an alkane or alkene series (C2 < C3 < C4 …).
Pattern 2 — anomalously high boiling points (NH₃, H₂O, HF): NH3, H2O and HF boil far higher than their group neighbours (e.g. PH3, H2S, HCl), even though they are smaller.
The reason is hydrogen bonding — the strongest IMF — because each has H bonded to N, O or F. PH3 has only weak London/dipole–dipole forces, so it boils far lower than NH3.
Worked example — comparing two molecules
Explain why water, H2O, boils at 100 °C but hydrogen sulfide, H2S, boils at −60 °C, even though H2S has more electrons.
Solution
- Identify the strongest force in each. Water has H bonded to O, so its molecules form hydrogen bonds. H2S has no H–N/O/F bond, so it has only weaker dipole–dipole and London forces.
- Compare the strength. Hydrogen bonds are much stronger than the forces in H2S.
- Link to boiling point. More energy is needed to separate the water molecules, so water boils at a much higher temperature.
Final answer
Water has hydrogen bonding (H on O); H2S has only dipole–dipole/London forces. The stronger hydrogen bonds in water need more energy to break, so water boils higher.
How this is tested: S2.2.5 is one of the most reliable Paper 2 explain questions, plus a Paper 1A 'name/identify the force' MCQ.
The standard ask is 'explain the boiling-point trend / difference in terms of intermolecular forces'. To score, you must: (1) name the IMF in each substance, (2) say which is stronger (and why — e.g. more electrons, or H–N/O/F present), and (3) link that to the energy needed and the boiling point.
The mark students lose: Never say boiling 'breaks the covalent bonds' or 'breaks the molecules apart'. Boiling only separates the molecules by overcoming the intermolecular forces — the covalent bonds stay intact. Saying 'breaks bonds' loses the mark.
IB-style question — the alkenes (a)
(a) The boiling points of the first four alkenes increase from ethene to but-1-ene. Explain this trend in terms of intermolecular forces. [2]
How to score the marks
- Mark 1 — identify the force and what changes. Alkenes are essentially non-polar, so the only intermolecular force is London (dispersion). As the chain lengthens, each molecule has more electrons (a larger, more polarisable molecule).
- Mark 2 — link to boiling point. More electrons → stronger London forces → more energy is needed to separate the molecules → the boiling point rises.
Final answer
London (dispersion) forces increase because larger alkenes have more electrons; stronger forces need more energy to overcome, so the boiling point rises along the series.
IB-style question — ammonia vs phosphine (b)
(b) Ammonia, NH3, boils at −33 °C, but phosphine, PH3, boils at −88 °C. Name the intermolecular force responsible for ammonia's much higher boiling point and explain your answer. [2]
How to score the marks
- Mark 1 — name the force. Ammonia molecules form hydrogen bonds, because hydrogen is bonded directly to nitrogen (one of N, O, F).
- Mark 2 — explain the difference. Phosphine has only weaker dipole–dipole and London forces (no H–N/O/F), so less energy is needed to separate PH3 molecules — hence ammonia boils much higher.
Final answer
Hydrogen bonding. NH3 has H bonded to N, so it hydrogen bonds; PH3 has only weaker dipole–dipole/London forces, so NH3 needs more energy to boil.