London Dispersion Forces: The Weakest Force That Matters Most

Why can nitrogen and oxygen be liquefied at all, if their molecules are non-polar and carry no charge? Why is iodine a solid at room temperature while fluorine is a gas? How does a gecko walk up a glass window? The answer to all three is the same: London dispersion forces. They’re the weakest of the intermolecular forces for small molecules, yet they act between every pair of molecules and atoms in existence, and for large molecules they can outweigh everything else. This article explains where they come from and what controls their strength.

A force between non-polar molecules

Most intermolecular forces arise from charges: ions, or permanent partial charges in polar molecules (see intermolecular forces). But noble gases such as argon and non-polar molecules such as methane, N₂ and I₂ have no permanent charge separation at all. Something must still attract them to each other, because they condense into liquids and freeze into solids when cooled enough. Argon liquefies at −186 °C; iodine is a solid at room temperature.

In 1930, the physicist Fritz London explained this attraction using quantum mechanics. The forces are named after him and are also called dispersion forces, instantaneous dipole–induced dipole forces, or (together with other weak forces) part of the family of van der Waals forces (see van der Waals forces).

How they arise: temporary dipoles

The electrons in an atom or molecule aren’t fixed in place. They’re constantly moving, and at any instant their distribution can be slightly uneven — a little more electron density on one side than the other. For that instant, the molecule has a temporary (instantaneous) dipole: one end slightly negative (δ−), the other slightly positive (δ+).

That temporary dipole affects its neighbours. Its δ− end repels the electrons of a nearby molecule, pushing them away and creating an induced dipole in the neighbour, lined up so that the two attract.

1. Electrons move randomly, creating a temporary dipole in molecule A.
2. The temporary dipole induces a dipole in neighbouring molecule B.
3. The δ+ end of one attracts the δ− end of the other.
4. A moment later the electrons shift, the dipoles change direction, and the process repeats — but the dipoles stay coordinated, so the attraction persists on average.

The attraction at any instant is tiny and fleeting, but it’s happening constantly between all molecules, and averaged over time it adds up to a real, measurable attraction.

What makes London forces stronger?

1. More electrons (larger molecules)

The more electrons a molecule has, and the further they are from the nuclei, the more easily its electron cloud can be distorted. This ease of distortion is called polarisability. More polarisable molecules form larger temporary dipoles and induce larger dipoles in their neighbours, so the London forces are stronger.

The halogens show this perfectly:

Halogen	Electrons per molecule	Boiling point	State at 20 °C
F₂	18	−188 °C	Gas
Cl₂	34	−34 °C	Gas
Br₂	70	59 °C	Liquid
I₂	106	184 °C	Solid

All four are non-polar diatomic molecules, so London forces are the only intermolecular forces. As the number of electrons rises, boiling points climb steeply — from a gas to a solid down one group. The noble gases follow the same pattern: helium boils at −269 °C, xenon at −108 °C.

2. Larger surface area (shape)

London forces act only over very short distances, so how much of each molecule can touch its neighbours matters. Molecules with the same number of electrons but different shapes can have quite different boiling points.

Pentane and 2,2-dimethylpropane both have the formula C₅H₁₂ and 42 electrons:

• Pentane is a long, straight chain. Molecules can lie alongside each other with lots of contact. Boiling point: 36 °C.
• 2,2-Dimethylpropane is compact and roughly spherical, like a ball. Molecules touch only at small points. Boiling point: 10 °C.

This is why branched alkanes have lower boiling points than their straight-chain isomers — a fact used in blending petrol.

3. Chain length

In a homologous series, each extra CH₂ group adds electrons and surface area, so boiling points rise steadily: methane (−162 °C), ethane (−89 °C), propane (−42 °C), butane (−1 °C), pentane (36 °C), hexane (69 °C), and so on. Long-chain alkanes are waxy solids. Polymers such as poly(ethene) have huge molecules held together almost entirely by London forces, which is enough to make tough plastics.

How strong are they?

For small molecules, London forces are weak — typically a few kJ mol⁻¹. But because they grow with size and contact area, they can become very large. In fact, for most molecules, London forces make up the largest share of the total intermolecular attraction, even for polar molecules. For hydrogen chloride, London forces contribute more to the total attraction than its permanent dipole does.

This explains a result that surprises many students: the boiling points of the hydrogen halides from HCl to HI increase (HCl −85 °C, HBr −67 °C, HI −35 °C), even though the molecules become less polar. The rising number of electrons, and so stronger London forces, outweighs the falling dipole–dipole attraction (see dipole–dipole forces). Hydrogen fluoride is the exception, with a boiling point of 20 °C, because it forms hydrogen bonds.

London forces in the real world

Geckos

A gecko can hang from a glass ceiling by one toe. Its feet are covered in millions of tiny hairs (setae), each splitting into hundreds of even tinier tips (spatulae), only about 200 nm across. These bring the gecko’s foot into extremely close contact with the surface over a huge total area. London forces between the tips and the surface — individually tiny — add up to enough to hold the animal. Research published in the early 2000s showed that the adhesion works on both water-attracting and water-repelling surfaces, pointing to van der Waals forces rather than capillary effects as the main mechanism.

Graphite and layered materials

In graphite, strong covalent bonds hold carbon atoms together within each layer, but only London forces hold the layers to each other. That’s why layers slide easily and why graphite rubs off onto paper from a pencil (see giant covalent structures).

Proteins and membranes

London forces between non-polar side chains help hold proteins in their folded shapes, and between the fatty tails of lipids they help hold cell membranes together (see protein folding).

Non-stick surfaces

PTFE (Teflon) is non-stick partly because fluorine atoms hold their electrons very tightly — they’re not very polarisable — so London forces between PTFE and food are unusually weak.

Common misconceptions

• “Only non-polar molecules have London forces.” All molecules and atoms have them; polar molecules have dipole–dipole forces in addition.
• “London forces are always the weakest.” For large molecules, they can be stronger than dipole–dipole forces or even hydrogen bonds in small molecules.
• “Heavier molecules boil higher because they’re heavier.” Mass correlates with electron count, but it’s the electrons and polarisability that matter, not mass itself.
• “London forces are permanent dipoles.” They come from temporary, constantly changing dipoles.

Key takeaways

• London dispersion forces are attractions between temporary and induced dipoles caused by moving electrons.
• They act between all atoms and molecules.
• They’re stronger for molecules with more electrons (greater polarisability) and larger contact area.
• They explain rising boiling points down group 17 and along homologous series, and why branched isomers boil lower.
• For large molecules, London forces are often the dominant intermolecular force.