Molecules are held together by covalent bonds. But what holds one molecule to the next one? Much weaker attractions called intermolecular forces — and they are the reason water is a liquid while carbon dioxide is a gas.
📚 What you need to know
Intramolecular = inside a molecule (covalent bonds, strong). Intermolecular = between molecules (weak).
There are four types: London (dispersion) forces, dipole–dipole, dipole–induced dipole, and hydrogen bonding.
Van der Waals forces is the umbrella term for the first three.
London forces act between all particles and get stronger with more electrons and bigger surface area.
Hydrogen bonding is the strongest and needs H bonded to N, O or F plus a lone pair on a neighbouring N, O or F.
Strength order: London < dipole–dipole < hydrogen bonding — and all are far weaker than a covalent bond.
Two different kinds of force
This distinction matters more than anything else on this page. When you boil water you get steam, which is still H2O — the molecules survive completely intact. What broke was the attraction between the molecules, not the O–H bonds inside them.
The orange arrow points at a covalent bond inside a molecule. The green dashes are the force between two separate molecules — that is what boiling has to overcome.
If someone asks why water boils at 100 °C, never say “the covalent bonds break”. They don’t. Boiling separates molecules from each other; it does not take them apart.
1. London (dispersion) forces
These exist between every pair of particles — atoms, polar molecules, nonpolar molecules, everything. They are also the only intermolecular force acting in nonpolar substances.
Here is how they arise. Electrons are constantly moving, so at any instant they might happen to be bunched slightly more on one side of a particle than the other. That instant of unevenness is a temporary dipole. It only lasts a moment, but while it exists it pushes the electrons in a neighbouring particle away, creating an induced dipole next door. Now the two particles have opposite ends facing each other, and they attract.
A temporary dipole appears by chance, induces one next door, and the two attract. This is happening constantly in every substance.
What makes London forces stronger
Two things, and both come up constantly in exam questions:
Number of electrons. More electrons means a bigger electron cloud, which is easier to distort, so the temporary dipoles are larger and the attraction is stronger. This is why boiling points rise as you go down a group.
Surface area / shape. Long, straight molecules can lie alongside each other and touch at many points. Branched, compact molecules touch at fewer points, so the forces are weaker.
Going down group 18 the atoms gain electrons, London forces strengthen, and the boiling point climbs steadily.
The shape effect in numbers.Pentane and 2,2-dimethylpropane are isomers — same formula C5H12, same number of electrons, so the same “raw” London force per contact point. But pentane is a long chain that packs closely alongside its neighbours, while 2,2-dimethylpropane is a compact ball that touches much less. Pentane boils at 36 °C; 2,2-dimethylpropane boils at 10 °C.
2. Dipole–dipole attractions
If a molecule is polar — it has a permanent dipole from the last set of notes — then it has an extra trick available. The δ− end of one molecule is attracted to the δ+ end of the next.
This is in addition to London forces, not instead of them. So a polar molecule always has stronger overall intermolecular forces than a nonpolar molecule of similar size.
Permanent dipoles line up head to tail. This attraction sits on top of the London forces that are already there.
WORKED EXAMPLE
Butane (C4H10) boils at 0 °C but propanone (CH3COCH3) boils at 56 °C. Both have the same number of electrons. Explain the difference.
Same number of electrons → similar London forcesSo London forces cannot be the reason.Butane is nonpolar — London forces onlyPropanone has a polar C=O bond and an asymmetric shape → permanent dipoleSo propanone has dipole–dipole attractions AS WELL AS London forces.more energy needed to separate propanone molecules → higher boiling point
3. Dipole–induced dipole attractions
This one is the hybrid case: a polar molecule sitting next to a nonpolar one. The permanent dipole of the polar molecule distorts the electron cloud of the nonpolar one, inducing a dipole in it, and the two then attract.
A good example is HCl mixed with Cl2. It is also the reason a small amount of oxygen will dissolve in water — which is what fish depend on.
4. Hydrogen bonding
This is the strongest intermolecular force, and it is really just a very extreme case of dipole–dipole attraction. It needs two ingredients, both of which must be present:
A hydrogen atom covalently bonded to N, O or F. These are so electronegative that they strip the hydrogen almost bare, leaving it very strongly δ+.
A lone pair on an N, O or F atom in a neighbouring molecule for that δ+ hydrogen to be attracted to.
Remember it as “N, O, F — and only these”. H bonded to carbon or chlorine is nowhere near polar enough. If a molecule has no N–H, O–H or F–H bond, it cannot hydrogen bond to itself, full stop.
Each water molecule can donate two hydrogen bonds through its O–H bonds and accept two more through its lone pairs.
How many hydrogen bonds a molecule can form depends on how many δ+ hydrogens it has and how many lone pairs it can offer:
Water has 2 O–H bonds and 2 lone pairs, so each molecule can form up to 4 hydrogen bonds. That is why water has such an unusually high boiling point for its size.
Ammonia has 3 N–H bonds but only 1 lone pair, so on average each molecule manages only 1 — which is why ammonia boils much lower than water.
Putting them in order
The order is what you need. Even hydrogen bonding is much weaker than the covalent bond it sits next to.
🧩 Working out which forces a substance has
Every substance has London forces. Start there and note the number of electrons.
Is the molecule polar? If yes, add dipole–dipole attractions.
Does it have H bonded directly to N, O or F? If yes, add hydrogen bonding.
To compare two substances, find the strongest force each one has — that usually decides it. Only fall back on electron count when both have the same type.
WORKED EXAMPLE
State all the intermolecular forces present in (a) Br2, (b) HCl, (c) CH3OH.
(a) Br₂ — identical atoms, so nonpolarLondon forces only(b) HCl — polar bond, linear so no cancellingLondon + dipole–dipoleNo hydrogen bonding: Cl is not N, O or F.(c) CH₃OH — has an O–H bondLondon + dipole–dipole + hydrogen bonding
💡 Exam tip
Use the full name London (dispersion) forces. “Induced dipole” or “instantaneous dipole” on its own may not get the mark.
Say “more electrons”, not “bigger molecule” or “heavier”. The electron count is what actually matters.
When comparing boiling points, name the type of force and say which is stronger and why. Two marks, two ideas.
Hydrogen bonding needs both parts — a δ+ H on N/O/F and a lone pair to aim at.
⚠️ Common mix-up
A hydrogen bond is not a covalent bond. It is an intermolecular attraction, roughly a twentieth of the strength.
Boiling breaks intermolecular forces, not covalent bonds. Steam is still H2O.
Polar molecules still have London forces. Dipole–dipole is an extra, not a replacement.
HCl does not hydrogen bond to itself. Chlorine is electronegative, but it is not N, O or F.
Van der Waals is a group name, not a fourth type of force. If a question asks which force, be specific.
Up next: Properties of Covalent Substances — putting these forces to work to explain melting points, volatility, solubility and conductivity.
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