London Dispersion Forces: Attraction in Every Molecule
How fleeting, random electron shifts create temporary dipoles that attract all molecules, and why bigger molecules boil at higher temperatures.
Quick answer
London dispersion forces are weak attractions present between all molecules, caused by momentary uneven electron distributions that create temporary dipoles and induce matching dipoles in neighbors.
Example: Iodine is a solid at room temperature while fluorine is a gas because iodine's large, easily distorted electron cloud creates far stronger temporary-dipole attractions between I2 molecules.
What you will learn
- Explain how instantaneous and induced dipoles create attraction.
- Use the halogen boiling-point trend as evidence for dispersion forces.
- Define polarizability and connect it to molecular size and shape.
- Show that dispersion forces act in polar molecules too.
Temporary dipoles from restless electrons
Electrons never sit still. At any instant, a molecule's electron cloud may bunch slightly to one side, giving a fleeting excess of negative charge there and a deficit on the other — an instantaneous dipole, living for a tiny fraction of a second.
That flicker does not go unnoticed. Its electric field nudges the electron cloud of a neighboring molecule, inducing a matching dipole: positive near the flicker's negative side, negative near its positive side. The two temporary dipoles attract — then the electrons shift again, and a new flicker takes over. Averaged over time, the attraction never cancels to zero.
Fritz London explained this quantum-mechanically in 1930, and the forces bear his name. The key insight: no permanent dipole is needed. Every molecule's electrons flicker, so every molecule — polar or nonpolar, noble gas or hydrocarbon — feels dispersion attraction. It is the universal intermolecular force.
Worked trend: boiling points of the halogens
The halogens provide the textbook demonstration. Fluorine (F2) boils at −188 °C, chlorine (Cl2) at −34 °C, bromine (Br2) at 59 °C, and iodine (I2) at 184 °C — a climb of over 370 degrees across the group.
All four are nonpolar diatomics, so dipole-dipole forces and hydrogen bonding are absent everywhere. The only attraction that can vary is London dispersion — and it varies enormously, because electron clouds grow larger and more distortable down the group.
Walk the logic: same molecular shape, same force type, steadily increasing size → steadily stronger attraction → steadily higher boiling point. The trend isolates dispersion as cleanly as any experiment in chemistry. When a single variable changes and the property follows monotonically, you have found the cause — and here the cause is the growing, sloshing electron cloud.
Polarizability: the master variable
The master variable behind dispersion strength is polarizability: how easily an electron cloud distorts. Large atoms with many electrons and diffuse outer shells polarize readily; small atoms with tight shells resist.
Electron count is the main lever — more electrons, more flicker, stronger induced dipoles. But shape matters too, as the C5H12 isomers prove. n-Pentane stretches out like a rod (bp 36 °C); neopentane balls up compactly (bp 9.5 °C). Same formula, same mass — yet the extended isomer boils 26 degrees higher because its larger surface area gives neighbors more contact for dispersion interactions.
This is why long, floppy molecules boil higher than compact ones of equal mass, and why branched alkanes are the low-boiling members of each family. Surface contact is dispersion's currency, and extended shapes spend it best.
London forces live in polar molecules too
A common error is filing dispersion under 'nonpolar molecules only.' In reality it operates everywhere — polar molecules flicker just like nonpolar ones, so they feel dispersion plus their dipole-dipole attractions.
Often dispersion contributes more than the dipole forces do. Recall the hydrogen halides: HI has the smallest dipole moment yet the highest boiling point, because iodine's dispersion attraction dwarfs the dipole differences. In large molecules generally, dispersion is the heavyweight and dipole-dipole the undercard.
Nature exploits this at scale. A gecko's foot carries millions of microscopic setae, each splitting into hundreds of spatulae — together creating enormous surface contact with the wall. The summed van der Waals attraction, mostly dispersion, supports the animal's weight on smooth glass. No glue, no suction: just temporary dipoles, multiplied millions of times.
From helium to candle wax
The noble gases trace dispersion's signature from the bottom up: helium boils at −269 °C, neon at −246 °C, argon at −186 °C, krypton at −153 °C, xenon at −108 °C. With no bonds, no dipoles, and no chemistry at all, the steady rise is pure polarizability — each heavier atom's cloud distorts more easily.
Hydrocarbons tell the same story at room temperature. Methane and ethane are gases; pentane is a liquid; paraffin wax — chains of 20+ carbons — is a solid. Each added CH2 group contributes more flickering electrons and more surface contact until dispersion alone locks the molecules into a solid.
Macroscopically, this is why petroleum jelly is thick, candle wax holds its shape, and helium barely liquefies even near absolute zero. One force, present in every molecule, scaling with size — from the coldest liquid known to the wax on your table.
Common mistakes
London dispersion acts between all molecules. It is the only force holding liquid nitrogen together and the reason iodine is a solid at room temperature.
For small molecules, yes — but in large ones they dominate. Iodine is a solid and long alkanes are waxes purely through dispersion attractions.
A single flicker is tiny, but summed over countless electrons and molecules the effect liquefies gases, holds wax together, and lets geckos climb glass.
Real-world applications
Explaining how noble gases liquefy at low temperatures
Designing lubricants, waxes, and greases
Predicting chromatography retention times
Developing gecko-inspired dry adhesives
Frequently asked questions
What causes London dispersion forces?
Random electron motion momentarily unbalances a molecule's charge cloud, creating an instantaneous dipole that induces a matching dipole in neighbors — a synchronized flicker of attraction, explained quantum-mechanically by Fritz London in 1930.
Why do bigger molecules boil higher?
Larger electron clouds are more polarizable: they form stronger temporary dipoles, so more thermal energy is needed to separate the molecules. Compare F2 (−188 °C) with I2 (184 °C).
Do polar molecules have dispersion forces?
Yes, always. Every molecule's electrons flicker; polar molecules feel dispersion plus dipole-dipole, and in big molecules dispersion usually contributes more to the total attraction.
What is polarizability?
How easily an electron cloud distorts. Large, diffuse clouds (iodine, long chains) polarize easily and attract strongly; small, tight clouds (helium, fluoride) barely do — which is why helium is hardest to liquefy.