Dipole-Dipole vs London Dispersion Forces: The Difference

Here is the question that trips up almost everyone the first time: helium is a single, perfectly symmetrical atom with no charges anywhere on it — so why does it ever turn into a liquid? If nothing pulls helium atoms together, liquid helium should not exist. It does. Something is attracting them.

The short answer: dipole-dipole forces are attractions between molecules that have a permanent positive end and negative end, so only polar molecules have them. London dispersion forces come from temporary, flickering imbalances in an electron cloud, so every atom and molecule has them — polar or not. Dispersion is the universal one; dipole-dipole is the bonus that polar molecules get on top.

Quick comparison at a glance

Feature	Dipole-dipole forces	London dispersion forces
What causes it	A permanent dipole — one end is always δ+, the other always δ−	A momentary, random pile-up of electrons on one side, which induces a matching dipole in a neighbour
Which particles have it	Polar molecules only	Everything — atoms, nonpolar molecules, polar molecules too
How long it lasts	Permanent, built into the molecule	Fleeting, appearing and vanishing constantly
Gets stronger when	The molecule is more polar (bigger dipole moment)	The particle has more electrons and is easier to distort, or has a flatter, more contactable shape
Rough energy range	About 5–25 kJ/mol	About 0.05–40 kJ/mol — can be tiny or very large
Everyday example	HCl, acetone, SO₂	Helium, Cl₂, petrol, candle wax

Both belong to the family called van der Waals forces, and both are far weaker than the covalent bonds inside a molecule — a distinction worth keeping straight.

What are dipole-dipole forces?

When two atoms of different electronegativity bond, the shared electrons sit closer to the greedier atom. That gives one end of the bond a partial negative charge (δ−) and the other a partial positive charge (δ+). If the molecule's shape doesn't cancel those out, the whole molecule ends up with a permanent dipole.

Now put a crowd of these molecules together. They don't tumble randomly for long — the δ+ end of one drifts toward the δ− end of another, like a box of tiny bar magnets shuffling into alignment. That electrostatic attraction is the dipole-dipole force.

Acetone (propanone) is a clean example. Its C=O bond is strongly polar, the oxygen carries a solid δ−, and acetone molecules stack head-to-tail in the liquid. That's why acetone boils at 56 °C while butane — almost the same molar mass, 58 g/mol, but nonpolar — is already a gas at room temperature, boiling at about −0.5 °C. Almost the same size, almost the same number of electrons (32 versus 34): the difference is almost entirely dipole-dipole.

What are London dispersion forces?

Electrons don't sit still. At any instant, a helium atom's two electrons might both happen to be on the left-hand side, leaving the right slightly bare. For that flicker of a moment, the atom is a dipole. And that momentary dipole tugs on the electrons of the atom next door, inducing a dipole there too — δ− facing δ+. The two attract. A heartbeat later the arrangement collapses and reforms somewhere else.

Average all those flickers over billions of particles and you get a real, continuous attraction: the London dispersion force, named after Fritz London, who explained it in 1930.

Two things control its strength:

• Number of electrons (polarizability). A big, electron-rich cloud is loose and squishy, so it distorts easily and produces bigger temporary dipoles. Watch the noble gases: helium boils at −269 °C, neon at −246 °C, argon at −186 °C, krypton at −153 °C, xenon at −108 °C. Nothing changes except how many electrons there are.
• Shape. Molecules that can lie flat against each other touch over more area, so they attract more. Pentane and neopentane are both C₅H₁₂, but pentane is a long chain that stacks like uncooked spaghetti and boils at 36 °C, while neopentane is a compact little ball that boils at 9.5 °C.

Which one is stronger?

This is where the standard classroom shortcut — "dispersion is the weakest force" — quietly falls apart.

For molecules of similar size, yes: adding a permanent dipole adds attraction, so the polar one boils higher. Acetone versus butane, above, is exactly that.

But dispersion scales hard with electron count, and it will happily overtake a modest dipole. Compare the hydrogen halides:

• HCl — clearly polar, dipole moment 1.08 D, 36.5 g/mol → boils at −85 °C
• HI — barely polar, dipole moment 0.45 D, 127.9 g/mol → boils at −35 °C

HI is the less polar molecule and it boils 50 degrees higher, because it has three times the electrons (54 versus 18) and its dispersion forces swamp HCl's dipole-dipole advantage. Iodine (I₂) is a solid at room temperature and has no dipole at all.

The honest rule: compare electron count first, then use polarity as the tie-breaker.

Worked examples

Try each before reading on.

1. Which boils higher, F₂ or Cl₂? Both nonpolar, so it's dispersion only — and chlorine has more electrons. Cl₂ wins (−34 °C versus −188 °C).

2. Which boils higher, propane (C₃H₈) or dimethyl ether (CH₃OCH₃)? Count electrons: both have exactly 26, so dispersion is a tie and polarity is the tie-breaker. Dimethyl ether is bent at its oxygen and polar; propane is not. Dimethyl ether boils at −24 °C, propane at −42 °C.

3. Which has stronger forces overall, Br₂ or HCl? Tempting to say HCl because it's polar. But Br₂ has 70 electrons to HCl's 18, and Br₂ is a liquid at room temperature while HCl is a gas. Dispersion wins.

Common mistakes to avoid

• "Nonpolar molecules have no intermolecular forces." They always have dispersion. If they didn't, nothing nonpolar could ever be liquefied — and petrol, oil, and liquid nitrogen all exist.
• "Dispersion is always the weakest." Only for small particles. In large molecules it is usually the dominant force, and it is the main reason candle wax is a solid.
• "Dipole-dipole and hydrogen bonding are the same thing." Hydrogen bonding is an unusually strong special case, which in practice means H bonded directly to N, O, or F. Ordinary dipole-dipole is much weaker.

FAQ

Is London dispersion a van der Waals force?
Yes. "Van der Waals forces" is the umbrella term covering dipole-dipole, dipole-induced dipole, and London dispersion forces. Hydrogen bonding is usually treated separately because it is so much stronger.

Do polar molecules have dispersion forces too?
Yes — every particle with electrons has them. A polar molecule has dispersion plus dipole-dipole, which is why polarity raises the boiling point of molecules of similar size.

Why do bigger molecules have stronger dispersion forces?
More electrons, held further from the nucleus, mean a cloud that distorts more easily. Bigger temporary dipoles mean bigger induced dipoles and a stronger average attraction.

Are these forces stronger than covalent bonds?
No — not close. A typical covalent bond costs 150–1000 kJ/mol to break; these forces are usually under 40 kJ/mol. Boiling water separates the molecules without breaking a single O–H bond.

The takeaway

Dispersion forces are universal and scale with electron count and shape; dipole-dipole forces are an extra attraction that only polar molecules enjoy. For similar-sized molecules the polar one wins, but a big enough electron cloud beats a modest dipole every time — so count electrons first, then check polarity.

Build on this with Hydrogen Bonds vs Van der Waals Forces: The Difference, Polar vs Nonpolar Bonds: What's the Difference?, and Intramolecular vs Intermolecular Forces. Next in this series: [What Is a Dipole Moment?] and [Boiling Point vs Melting Point].

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