Intermolecular Forces
The weak attractions between molecules decide whether something is a gas, a liquid or a solid — and they are the reason water behaves like nothing else.
Builds on: 1.4 Molecular Shapes & Why Water Is Bent1.3 Metallic Bonds & the Bond Spectrum
The forces between, not within
Covalent bonds hold a molecule together. Intermolecular forces hold separate molecules to each other. They are far weaker — typically a few percent as strong — and they never appear in a chemical equation. But they decide almost every physical property you can observe: melting point, boiling point, viscosity, surface tension, whether something dissolves.
Three kinds matter, weakest first.
1. London dispersion forces (in everything)
Electrons move. At any instant they may happen to bunch on one side of a molecule, creating a fleeting dipole, which induces an opposite fleeting dipole in a neighbour — and for that moment the two attract. Averaged over time it is a real, if feeble, force.
Dispersion forces exist in every substance and grow with the number of electrons. That is why the noble gases liquefy in order down the group, and why methane is a gas while candle wax — the same kind of molecule, just far longer — is a solid.
2. Dipole–dipole forces (in polar molecules)
Polar molecules have permanent δ+ and δ− ends, so they line up and attract head to tail. Stronger than dispersion for molecules of similar size, which is why polar substances generally boil higher than nonpolar ones of comparable mass.
3. Hydrogen bonding (the strong one)
A special, unusually strong case of dipole–dipole. It requires hydrogen bonded directly to N, O or F — the three small, greedy atoms. Hydrogen has just one electron, so when oxygen or fluorine pulls that electron away, what remains is close to a bare proton: a very concentrated positive charge that can get very close to a lone pair on a neighbouring molecule.
Hydrogen bonds need H attached to Nitrogen, Oxygen or Fluorine. H–Cl does not qualify: chlorine is polar enough but too large, so its charge is too spread out. The German memory hook is short: „NOF — sonst nix.“
The anomaly that proves it
Compare the hydrides of group 16 — H₂O, H₂S, H₂Se, H₂Te. Dispersion forces grow with size, so boiling points should rise steadily down the group. They do, for the last three. Water, the smallest and lightest, should boil around −80 °C.
It boils at +100 °C. That 180-degree discrepancy is hydrogen bonding, and it is why there is liquid water on this planet at all.
Water is exceptional even among hydrogen-bonded substances, because of its shape. Each molecule has two hydrogens to donate and two lone pairs to accept — so every water molecule can hold four neighbours at once, building a continuous network. Ammonia has three hydrogens but only one lone pair; hydrogen fluoride has one hydrogen and three lone pairs. Only water is balanced.
Follow the consequences:
- Ice floats. Freezing locks the network into an open hexagonal cage that is less dense than the liquid. Almost every other substance sinks in itself. Lakes therefore freeze from the top down, and life survives underneath.
- Huge heat capacity. Warming water means loosening a vast web of hydrogen bonds, which absorbs enormous energy — oceans moderate the climate, and sweat cools you efficiently.
- Surface tension. Molecules at the surface are pulled inwards by neighbours with nothing above to balance them. Insects walk on it.
- “Like dissolves like.” Polar water dissolves polar and ionic things and refuses nonpolar ones, because it would have to break its own hydrogen-bond network to make room for a molecule that offers nothing in return.
⚗️ Lab — Boiling Point Lab
Four small molecules, one thermometer. The purple lines are the attractions between molecules.
- Set the temperature to −50 °C and step through all four substances. Two are gases, two are liquids.
- Compare CH₄ (M = 16) with H₂O (M = 18) — nearly the same mass, 261 degrees apart in boiling point.
- Raise the temperature past the boiling point and watch the attraction lines let go.