This chemistry grades 11-12 lesson on intermolecular forces notes gives the key ideas first, then worked examples and diagrams. After that, practise with the worksheet on intermolecular forces and check yourself with the quiz.
Quick answer: Intermolecular forces are the attractions between molecules: London forces act between all particles and grow with the number of electrons, permanent dipole–dipole forces act between polar molecules, and hydrogen bonds act when H bonded to N, O or F meets a lone pair. Stronger forces mean higher boiling points.
Level: Chemistry grades 11-12 (A Level · IB · Class 11-12) · Curriculum: IB S2.2
Intermolecular forces notes: the key ideas
What are intermolecular forces?
Inside a molecule, atoms are held by strong covalent bonds. Between molecules there are much weaker attractions. When water boils or iodine sublimes, it is these weak forces that break, not the covalent bonds. A typical O–H covalent bond needs about 460 kJ mol⁻¹ to break, while the attraction between neighbouring water molecules is about 20 kJ per mole of hydrogen bonds.
Intermolecular forces are attractions between neighbouring molecules (or atoms, for noble gases). Their strength decides the melting point, the boiling point and the volatility of a molecular substance.
How do London forces arise?
Electrons are always moving, so at any instant the charge in an atom or molecule is slightly unevenly spread. This instantaneous dipole induces an opposite dipole in a neighbour, and the two attract. London (dispersion) forces exist between all particles, even non-polar ones.
London forces grow stronger as the number of electrons increases, so larger molecules have higher boiling points. For the same number of electrons, long straight chains pack more contact surface than compact, branched molecules, so they also boil higher.
Alkanes show the same pattern: methane boils at −162 °C, butane at −1 °C and hexane at 69 °C. Branching lowers the boiling point: pentane boils at 36 °C, but its branched isomer 2,2-dimethylpropane boils at 9.5 °C.
What are permanent dipole–dipole forces?
A molecule with a net dipole, such as HCl or propanone, has a permanent δ+ end and δ− end. Neighbouring molecules line up so that opposite ends attract. These forces come in addition to London forces.
Fluorine, F₂, and hydrogen chloride, HCl, each have 18 electrons. F₂ boils at −188 °C and HCl at −85 °C. Explain the difference.
The number of electrons is equal, so the London forces are similar.
HCl is polar, so it also has permanent dipole–dipole forces, which F₂ lacks. More energy is needed to separate HCl molecules, so HCl has the higher boiling point.
What is a hydrogen bond?
A hydrogen bond is a strong intermolecular attraction between a hydrogen atom bonded to N, O or F (δ+) and a lone pair on a different N, O or F atom (δ−). The three atoms lie close to a straight line.
Hydrogen bonds are stronger than ordinary dipole–dipole forces because the small, highly polarised H atom can approach the lone pair very closely. They explain why water (100 °C) boils far above H₂S (−60 °C), as the graph shows.
Ethanol (C₂H₅OH) boils at 78 °C and its isomer methoxymethane (CH₃OCH₃) at −24 °C. Explain.
Both have 26 electrons, so London forces are similar.
Ethanol has an O–H group, so molecules form hydrogen bonds. Methoxymethane has no H bonded to O, so it has only dipole–dipole and London forces. Ethanol therefore boils far higher.
Boiling water does not break the O–H bonds. Writing that “covalent bonds break when water boils” is wrong: only the intermolecular hydrogen bonds are broken, and the molecules stay as H₂O.
| Force | Present in | Relative strength |
|---|---|---|
| London forces | all particles | weakest, grows with electrons |
| Permanent dipole–dipole | polar molecules such as HCl, propanone | stronger than London for equal size |
| Hydrogen bonding | molecules with N–H, O–H or F–H | strongest intermolecular force |
To rank boiling points, go through three steps: first check for hydrogen bonding, then for a permanent dipole, and finally compare the number of electrons. Compare molecules with similar numbers of electrons first.
- London forces: instantaneous dipoles, stronger with more electrons
- Permanent dipole–dipole forces act between polar molecules
- Hydrogen bond: H on N, O or F attracted to a lone pair on N, O or F
- Boiling breaks intermolecular forces, not covalent bonds
- Branching lowers boiling points by reducing surface contact
London, dipole and hydrogen bond attractions between molecules: frequently asked questions
Why does the boiling point increase down a group of noble gases?
Each heavier noble gas has more electrons, so its electron cloud is larger and more easily distorted. London forces become stronger, and more energy is needed to separate the atoms.
What conditions are needed for hydrogen bonding?
A hydrogen atom must be bonded to nitrogen, oxygen or fluorine, and a different molecule must have a lone pair on N, O or F. The H atom then points directly at that lone pair.
Why does branching lower the boiling point?
A branched molecule is more compact, so it touches its neighbours over a smaller surface. The London forces are weaker, and less energy is needed to separate the molecules.
Are intermolecular forces stronger than covalent bonds?
No. Covalent bonds are typically 10 to 100 times stronger. That is why changes of state, which only overcome intermolecular forces, need much less energy than chemical reactions that break bonds.
Why does ice float on water?
In ice each molecule forms four hydrogen bonds in an open hexagonal lattice with empty space. This makes ice about 8 to 9 % less dense than liquid water, so it floats.
