Which Substances Exhibit Only London Dispersion Forces

Some molecules hold hands. That said, others just sort of… brush up against each other. And if you've ever wondered why helium refuses to become a liquid until it gets absurdly cold, or why methane boils at a temperature where you'd never casually encounter it as a liquid — the answer lives in a single, deceptively simple type of intermolecular force: London dispersion forces.

These are the weakest of the major intermolecular attractions. But don't let "weakest" fool you into thinking they're unimportant. In the right molecules, they build up into something surprisingly powerful. And right now, I want to focus on a cleaner question: which substances exhibit only* London dispersion forces, and what that actually means for how they behave.

What Are London Dispersion Forces, Exactly?

Before we get into which substances are governed by them, let's make sure we're on the same page about what they actually are.

London dispersion forces (sometimes called London forces* or instantaneous dipole-induced dipole forces*) arise from the random movement of electrons within atoms and molecules. At any given moment, the electron cloud around an atom isn't perfectly symmetric. That tiny, temporary asymmetry creates a fleeting partial negative charge on one side and a fleeting partial positive charge on the other. That momentary dipole then induces a similar ripple in a neighboring atom or molecule, and the two get drawn toward each other — just for an instant.

It happens everywhere, all the time. Every substance has London dispersion forces at play. Even water, which is held together by strong hydrogen bonds, has LDF contributing to the overall picture.

The key thing is that LDF are always present*. But some substances have additional* intermolecular forces layered on top — permanent dipole-dipole interactions, hydrogen bonding, ionic interactions. Those other forces are stronger, so they tend to dominate the behavior of the substance.

A substance that exhibits only* London dispersion forces is one where every other type of intermolecular attraction is essentially absent. No permanent dipole. Which means no hydrogen bonding. No ionic character.

Why "Only" Is a More Specific Claim Than It Sounds

Here's where I want to be careful, because this is a nuance that trips a lot of people up. So technically, every* intermolecular interaction ultimately traces back to electromagnetic forces, which includes London forces. So saying a substance has "only" LDF means it lacks dipole-dipole interactions, hydrogen bonding, and ionic bonding — not that LDF are the only electromagnetic phenomenon in existence.

The practical effect is that these substances tend to have very low boiling and melting points, because there's nothing much holding the molecules together once you give them a little energy. Because of that, methane boils at −161°C. Helium boils at −269°C. Neon, nitrogen, argon — all famously reluctant to become liquids.

Which Substances Have Only London Dispersion Forces?

Here's the core of the question. The substances that exhibit only* London dispersion forces fall into a few clean categories.

Noble Gases

This is the purest example. Helium, neon, argon, krypton, xenon, and radon are all monatomic — they exist as single atoms, not bonded to anything else. Because the atom itself is nonpolar and has no permanent dipole moment, the only* intermolecular force acting between noble gas atoms is London dispersion.

This is why noble gases have such dramatically low boiling points. Helium's boiling point is just 4.2 kelvin. There's essentially no electrostatic attraction pulling the atoms together, so you have to strip away almost all their thermal energy before they'll condense.

Diatomic and Small Nonpolar Molecules

Several diatomic molecules are nonpolar because both atoms are the same — meaning the electron cloud is shared equally and there's no permanent dipole. The major players here:

• Nitrogen (N₂) — Makes up about 78% of the air you breathe. Under normal conditions it's a gas, and it liquefies at −196°C.
• Oxygen (O₂) — Also nonpolar, liquefies at −183°C.
• Hydrogen (H₂) — The smallest diatomic molecule, boils at −253°C.
• Chlorine (Cl₂) — Gas at room temperature, liquefies at −34°C, so you can actually observe it condensing under pressure.
• Fluorine (F₂) — Boils at −188°C.

These molecules have no dipole moment, no hydrogen attached to an electronegative atom, and no ionic character. The only intermolecular attraction is LDF. That said, larger molecules in this category — like Br₂, which is liquid at room temperature — have significantly stronger LDF due to their larger electron clouds, even though the type* of force hasn't changed.

Carbon Dioxide and Other Small Nonpolar Molecules

Carbon dioxide (CO₂) is a great example because it looks* like it might have dipole-dipole interactions. The molecule is linear (O=C=O), and the two oxygen atoms pull electron density away from the carbon. But here's the thing: the symmetry of the molecule means those dipoles cancel out. CO₂ has no net dipole moment*. It's nonpolar, so it relies entirely on London dispersion forces.

That's why dry ice (solid CO₂) sublimes directly at −78°C rather than melting into a liquid at atmospheric pressure — there's simply not enough attraction between CO₂ molecules to hold them in a liquid phase under normal conditions.

Want to learn more? We recommend how many minutes are in 6 hours and what goes in the water black and comes out red for further reading.

Carbon monoxide (CO) is a bit of an edge case. It does have a small dipole

Carbon Monoxide – A Tiny Dipole, But Still LDF‑Dominated

Carbon monoxide (CO) possesses a small permanent dipole moment (≈ 0.This means CO’s boiling point (‑191.5 °C) is only a few degrees higher than that of nitrogen, reflecting the modest extra stabilization from its dipole. 112 D) because the electronegativity difference between carbon and oxygen is modest and the molecular orbital arrangement leaves a slight excess of negative charge on the carbon side. Even so, the dipole–dipole interaction energy is minuscule compared with the instantaneous‑dipole/induced‑dipole (London) contribution for such a light molecule. In practice, CO is still classified as a non‑ dipole‑dominant substance; the bulk of its intermolecular attraction stems from London dispersion forces.

Larger Non‑Polar Hydrocarbons

As the size and number of electrons in a molecule increase, polarizability rises sharply, and the strength of LDF grows even though the type* of force remains unchanged. This explains the monotonic increase in boiling points across a homologous series:

-butane (C₄H₁₀) | –0.5 | 34 | ★★★★

Notice how the boiling point climbs by roughly 40–80 °C with each additional CH₂ unit, even though every molecule remains perfectly non‑polar. The trend is a textbook illustration of how dispersion forces scale with electron count and surface area.

Fluorine, Chlorine, and the Halogens

The diatomic halogens sit neatly in the “LDF‑only” column of the intermolecular‑force table:

• F₂ – boiling point –188 °C; 18 electrons.
• Cl₂ – boiling point –34 °C; 34 electrons.
• Br₂ – boiling point 59 °C; 70 electrons; liquid at room temperature.
• I₂ – melting point 114 °C, boiling point 184 °C; 106 electrons; solid at room temperature.

Each jump in size brings a larger, more diffuse electron cloud that can be polarized more easily, strengthening the instantaneous‑dipole interactions. No dipole‑dipole or hydrogen‑bonding terms enter the picture because the X–X bond is perfectly symmetric and the atoms are identical.

Noble Gases – The Purest LDF Systems

If you want the cleanest example of pure London dispersion forces, look no further than the noble gases:

Helium boils barely above absolute zero, while xenon condenses at a temperature where water would be a solid. Yet every member of the group is monatomic, non‑polar, and lacks any permanent multipole moments. The entire range of physical behavior is dictated by the number* and diffuseness* of electrons available for temporary polarization.

Why This Matters in Everyday Chemistry

Understanding that non‑polar substances rely solely on London dispersion forces is more than a textbook exercise. It explains:

• Why oil and water don’t mix. Long‑chain hydrocarbons are held together by LDF; water is held together by strong hydrogen bonds. Mixing would require breaking H‑bonds without compensating LDF, which is energetically unfavorable.
• Why dry ice “disappears.” The sublimation of CO₂ at –78 °C is a direct consequence of weak LDF between linear, non‑polar molecules.
• Why larger molecules have higher boiling points. Even a single extra CH₂ group can add tens of degrees to a molecule’s boiling point, simply by enlarging the electron cloud.
• Why certain non‑polar gases must be stored at cryogenic temperatures. Helium and neon, with only a handful of electrons, require liquid‑helium temperatures to liquefy at all.

Closing Thoughts

London dispersion forces are the universal* intermolecular interaction—every molecule, regardless of polarity, experiences them. Also, for non‑polar species, they are the only game in town, and their magnitude depends on how many electrons are present and how easily those electrons can be distorted. Because of that, from the fleeting dipoles that hold argon atoms together at –186 °C to the substantial LDF network that keeps bromine a liquid at room temperature, the story of non‑polar substances is, at its heart, a story about electron clouds, polarizability, and the subtle, ever‑present push‑and‑pull of quantum fluctuations. Recognizing this allows chemists to predict phase behavior, solubility trends, and even the practical handling of everything from laboratory gases to industrial solvents.