Yes, every molecule has London dispersion forces. Every atom does too. These forces are universal because they arise from the movement of electrons, and every piece of matter has electrons. Whether a molecule is polar or nonpolar, small or large, simple or complex, London dispersion forces are always present and always contributing to how that substance behaves.
Why Every Molecule Has Them
London dispersion forces originate from temporary, correlated shifts in electron distribution. At any given instant, the electrons in an atom or molecule aren’t perfectly evenly spread. They might cluster slightly toward one side, creating a fleeting negative charge on that side and a fleeting positive charge on the other. This momentary imbalance is called an instantaneous dipole.
That temporary dipole then influences a neighboring atom or molecule. The slightly negative side repels the neighbor’s electrons, pushing them away and inducing a matching dipole in the neighbor. For a brief moment, the two particles are attracted to each other, positive end facing negative end. The dipoles vanish almost immediately, but new ones constantly form and re-form, creating a net attractive force over time.
Because this mechanism requires nothing more than electrons in motion, it applies to everything: nonpolar molecules like oxygen gas, polar molecules like water, noble gas atoms like argon, and massive structures like proteins. A molecule can have other intermolecular forces layered on top (dipole-dipole interactions, hydrogen bonds), but London dispersion forces are always part of the package.
Noble Gases: Proof the Force Is Universal
Noble gases are the clearest demonstration. Argon, helium, and neon are single atoms with no bonds and no permanent charge imbalance. Yet if you cool argon to about negative 186°C, it condenses into a liquid. Something must be holding those atoms together, and that something is London dispersion forces. They’re the only attractive force available.
The noble gases also show how these forces scale. Helium, the smallest noble gas with just two electrons, boils at negative 269°C, barely above absolute zero. Argon, which is larger and has more electrons, boils at negative 186°C. The pattern holds across the entire group: more electrons mean stronger dispersion forces, which means a higher boiling point.
What Makes Dispersion Forces Stronger or Weaker
Not all London dispersion forces are equal. Their strength depends mainly on how easily a molecule’s electron cloud can be distorted, a property called polarizability. Several factors influence this:
- Size and electron count. Larger atoms and molecules with more electrons have valence electrons that sit farther from the nucleus. Those outer electrons are held less tightly and shift more easily, creating stronger temporary dipoles. This is why iodine (a large molecule) is a solid at room temperature while fluorine (a small one) is a gas, even though both are nonpolar.
- Molecular shape. Long, stretched-out molecules have more surface area available for contact with neighboring molecules. More contact means more opportunities for temporary dipoles to form and interact. A straight-chain hydrocarbon, for example, has a higher boiling point than a highly branched version with the same number of atoms, because the linear shape allows molecules to line up closely.
How They Compare to Other Intermolecular Forces
For small molecules, London dispersion forces are typically the weakest intermolecular force. Dipole-dipole interactions in small polar molecules are significantly stronger, and hydrogen bonds (found in water, alcohols, and similar compounds) are stronger still. This is why many students assume dispersion forces are trivial.
But that assumption breaks down with larger molecules. Because dispersion forces scale with size and electron count, they can actually dominate in big, heavy species. Buckminsterfullerene (C₆₀), a nonpolar molecule made of 60 carbon atoms, has a boiling point above 280°C. That’s far higher than nitrous oxide (N₂O), a polar molecule that boils at negative 88.5°C. The sheer number of electrons in C₆₀ generates dispersion forces strong enough to overwhelm the dipole-dipole interactions of the smaller polar molecule.
The Common Misconception About Polar Molecules
A frequent mistake in chemistry courses is thinking that polar molecules have dipole-dipole forces instead of London dispersion forces. In reality, polar molecules have both. Water, for instance, experiences hydrogen bonding, dipole-dipole interactions, and London dispersion forces simultaneously. The hydrogen bonds happen to be so much stronger that they dominate water’s behavior, but the dispersion forces are still there, contributing to the total attraction between water molecules.
When listing the intermolecular forces present in a substance, dispersion forces should always be on the list. The question is never “does this molecule have London dispersion forces?” but rather “how significant are the London dispersion forces compared to whatever else is going on?”
Dispersion Forces in Biology
These seemingly weak forces play a surprisingly large role in biological systems. Proteins fold into specific three-dimensional shapes, and London dispersion forces are one of the key reasons they stay folded. When a protein folds, its nonpolar amino acid side chains get buried in the interior, packed tightly together. The tight packing enhances London dispersion forces between those buried groups, and research on protein stability has shown that these van der Waals interactions make an important contribution to keeping the structure intact.
The effect is cumulative. A single dispersion interaction between two small groups is negligible. But a protein has hundreds or thousands of atoms packed into its core, each participating in dispersion interactions with its neighbors. Added together, these tiny forces become a significant stabilizing factor. It’s a good example of how a force that seems trivial at the scale of two helium atoms becomes essential at the scale of a biological molecule with thousands of electrons.