Intermolecular forces

Intermolecular forces are the weak attractions between molecules. They are much weaker than the covalent bonds inside molecules, but they decide the melting and boiling points of simple molecular substances. There are three kinds: London forces (induced dipole–dipole forces), permanent dipole–dipole forces and hydrogen bonds.

Key facts

  • London forces act between all molecules: a temporary dipole in one molecule induces a dipole in its neighbour, and the two attract. More electrons make bigger temporary dipoles and stronger forces.
  • Long, straight molecules have stronger London forces than branched ones with the same number of electrons, because they touch over a larger surface: pentane boils at 36 °C, 2,2-dimethylpropane at 10 °C.
  • Permanent dipole–dipole forces act between polar molecules such as HCl and propanone: the δ+ end of one molecule attracts the δ− end of the next.
  • A hydrogen bond forms between a hydrogen atom bonded to nitrogen, oxygen or fluorine and a lone pair on a nitrogen, oxygen or fluorine atom of another molecule. Hydrogen bonds are the strongest intermolecular forces.
  • When a simple molecular substance melts or boils, only the intermolecular forces are overcome; the covalent bonds inside the molecules stay.
  • Hydrogen bonding explains why water boils at 100 °C (H₂S at −60 °C), why ice is less dense than water, and why alcohols boil at higher temperatures than alkanes of a similar size.

Worked examples

Example 1

Explain why the boiling points of the halogens increase from fluorine (−188 °C) to iodine (184 °C).

  1. Halogen molecules are non-polar, so only London forces act between them.
  2. Down the group the molecules get bigger: F₂ has 18 electrons, I₂ has 106.
  3. More electrons give bigger temporary dipoles, so the London forces are stronger.
  4. More energy is needed to separate the molecules, so the boiling point rises.

Answer: Down the group the molecules have more electrons, so the London forces between them are stronger and need more energy to overcome.

Example 2

Ethanol, C₂H₅OH, boils at 78 °C, but propane, C₃H₈, with almost the same number of electrons, boils at −42 °C. Explain the difference.

  1. Ethanol has an O–H group. Oxygen is very electronegative, so the hydrogen atom is strongly δ+.
  2. That hydrogen is attracted to a lone pair on the oxygen of another ethanol molecule: a hydrogen bond.
  3. Propane has no hydrogen atom on N, O or F, and is almost non-polar: only London forces, which are about the same size as ethanol’s.
  4. Hydrogen bonds are much stronger, so more energy is needed to boil ethanol.

Answer: Ethanol molecules form hydrogen bonds with each other; propane molecules have only London forces.

Example 3

Hydrogen chloride (−85 °C) is more polar than hydrogen iodide (−35 °C). Why does hydrogen iodide have the higher boiling point?

  1. Both are polar, so both have London forces and permanent dipole–dipole forces.
  2. HCl is more polar, so its dipole–dipole forces are a little stronger.
  3. But HI has 54 electrons and HCl only 18, so HI has much stronger London forces. 1 + 53 = 54; 1 + 17 = 18
  4. For these molecules London forces matter more, so HI boils at the higher temperature.

Answer: HI has many more electrons, so its London forces are much stronger; this outweighs HCl’s stronger dipole–dipole forces.