The forces with a Dutch surname
In the last guide we met the family name. The van der Waals forces are an umbrella term for the everyday attractions between molecules that don't carry a full electric charge. Johannes van der Waals first realized, while studying gases, that real molecules tug on each other a little — and that this small tug explains why gases can be squeezed into liquids. The umbrella covers three members; today we meet the most universal of them.
The three members are: dispersion forces (the subject of this guide), dipole–dipole interactions, and the special case of hydrogen bonding (both in the next guide). They differ in strength and in which molecules feel them — but all three are gentle, short-range tugs that vanish the moment molecules move apart.
A force from nothing: the flickering cloud
Here is the strange one. Take two molecules with no permanent lopsidedness at all — say, two argon atoms, perfectly symmetric. On average their electrons are spread evenly, so neither has a positive or negative end. And yet, cool argon enough and it becomes a liquid. Something pulls. What?
The answer is that *on average* even is not the same as even *at every instant*. Electrons are restless; they slosh around the molecule. At any one moment, slightly more of them might pile up on one side, leaving that side momentarily negative and the other side positive. For a flicker of an instant, the symmetric atom has a tiny temporary dipole — a fleeting positive end and negative end.
- In molecule A, electrons momentarily slosh to one side — A now has a fleeting negative end.
- A's negative end pushes the electrons of nearby molecule B away, so B's near side turns positive — A has *induced* a dipole in B.
- Now A's negative end faces B's positive end: opposite charges, so they attract for that instant.
- The flicker shifts, but a new one appears somewhere else. Averaged over countless flickers, the net effect is a steady, gentle pull.
This is the London dispersion force, named after the physicist Fritz London who explained it. It is the weakest intermolecular force — but it is also the only one *every* substance has, because every molecule has restless electrons. Even the noblest, most stand-offish atoms feel it.
Why bigger molecules stick harder
Dispersion forces are weak per molecule, but two things make them stronger. The first is size. A big molecule has many more electrons, spread over a larger, fluffier cloud. Such a cloud is easy to distort — its electrons slosh more freely, making bigger temporary dipoles. We call this squishiness polarizability: how readily a molecule's electron cloud deforms in an electric field.
More polarizable molecules have stronger dispersion forces, and that shows up directly in boiling points. The halogen molecules are a clean example: tiny fluorine and chlorine are gases at room temperature, bigger bromine is a liquid, and the largest, iodine, is a solid. Nothing changed but size — and with it, polarizability, and with that, the grip of dispersion.
The second thing is shape. Long, stretched-out molecules can lie alongside each other like logs in a stack, touching along their whole length, so the little tugs add up. The same molecule rolled into a compact ball touches its neighbours at only a few points. This is why straight-chain candle wax is more solid than a branched isomer of the same formula — same atoms, but more surface in contact.
Weak, but everywhere — and it adds up
It is tempting to dismiss dispersion as too feeble to matter. That would be a mistake. A single dispersion contact is weak, but a large molecule makes thousands of them at once, and thousands of feeble tugs become a firm hold. Candle wax, plastic, and the fats in your body are held together almost entirely by dispersion forces — and they are very much solid or sticky at room temperature.