Dipole–Dipole Forces Explained

Propane and ethanal (acetaldehyde) have almost the same molar mass — 44 g mol⁻¹ each. Yet propane boils at −42 °C, and ethanal at 20 °C. The difference is a type of intermolecular attraction that propane can’t form but ethanal can: permanent dipole–dipole forces. This article explains how these forces arise, how strong they are compared with other intermolecular forces, and how to use them to explain boiling points and other properties.

Starting point: polar molecules

A polar molecule has a permanent, uneven distribution of charge: one end is slightly negative (δ−) and the other slightly positive (δ+). This happens when a molecule contains polar bonds — bonds between atoms of different electronegativity — and its shape doesn’t allow the bond polarities to cancel (see bond polarity and how to tell if a molecule is polar).

Examples of polar molecules: HCl, H₂O, NH₃, CH₃Cl, propanone (CH₃COCH₃), ethanal (CH₃CHO). Examples of non-polar molecules: H₂, CO₂, CH₄, CCl₄, BF₃.

The size of the charge separation is measured by the dipole moment (see dipole moments).

What dipole–dipole forces are

Permanent dipole–dipole forces are the electrostatic attractions between the δ+ end of one polar molecule and the δ− end of a neighbouring polar molecule.

In a liquid or solid, polar molecules tend to line up, head to tail, so that δ+ ends sit near δ− ends:

δ+H–Clδ− ··· δ+H–Clδ− ··· δ+H–Clδ−

Molecules in a liquid are constantly moving and tumbling, so this alignment is never perfect. But on average, attractive arrangements are slightly favoured over repulsive ones, giving a net attraction. As temperature rises, the molecules move more energetically and the alignment becomes less effective, so dipole–dipole forces weaken with temperature.

How strong are they?

Dipole–dipole forces are typically in the range of about 5–25 kJ mol⁻¹, depending on the molecule. That’s:

• stronger than London forces between small molecules of similar size;
• weaker than hydrogen bonds, which are an especially strong type of dipole–dipole interaction (see hydrogen bonding);
• much weaker than covalent or ionic bonds.

A key point: polar molecules have London forces as well. Every molecule has London forces; polar molecules have dipole–dipole forces in addition. So the total attraction between polar molecules is the sum of both (see London dispersion forces).

Worked comparison 1: propane vs ethanal

	Propane, CH₃CH₂CH₃	Ethanal, CH₃CHO
Molar mass	44 g mol⁻¹	44 g mol⁻¹
Electrons	26	24
Polar?	Almost non-polar (μ ≈ 0.08 D)	Polar (μ ≈ 2.7 D) — C=O bond
Intermolecular forces	London only	London + dipole–dipole
Boiling point	−42 °C	20 °C

With similar size and electron count, their London forces are similar. Ethanal’s strongly polar C=O group adds dipole–dipole attractions, so more energy is needed to separate its molecules, and it boils 62 °C higher.

Worked comparison 2: cis vs trans isomers

The two isomers of 1,2-dichloroethene have the same formula (C₂H₂Cl₂), the same atoms and the same bonds:

• cis: both Cl atoms on the same side of the C=C. The C–Cl bond dipoles add up → polar (μ ≈ 1.9 D). Boiling point 60 °C.
• trans: Cl atoms on opposite sides. The bond dipoles cancel → non-polar. Boiling point 48 °C.

Same size, same London forces — the difference in boiling point comes from dipole–dipole forces in the cis isomer.

Worked comparison 3: when London forces win

It’s tempting to assume the more polar molecule always boils higher. Look at the hydrogen halides:

Molecule	Dipole moment (D)	Electrons	Boiling point
HCl	1.08	18	−85 °C
HBr	0.82	36	−67 °C
HI	0.44	54	−35 °C

Polarity decreases from HCl to HI, yet boiling point increases. The growing number of electrons makes London forces stronger, and that outweighs the weakening dipole–dipole forces. For HI, London forces account for almost all of the attraction.

Lesson: when comparing substances, consider both size (London forces) and polarity (dipole–dipole). Dipole–dipole forces are the deciding factor when molecules are similar in size.

Where dipole–dipole forces matter

Solubility

Polar molecules dissolve well in polar solvents, partly because solute and solvent molecules can attract each other through dipole–dipole forces. Propanone (acetone) mixes completely with water and is also a good solvent for many polar organic substances (see “like dissolves like”).

Boiling points of organic compounds

Aldehydes and ketones (containing C=O) boil higher than alkanes of similar size, but lower than alcohols, which can hydrogen-bond:

Compound	M (g mol⁻¹)	Main extra force	Boiling point
Butane	58	none (London only)	−1 °C
Propanone	58	dipole–dipole	56 °C
Propan-1-ol	60	hydrogen bonding	97 °C

Liquid crystals

Liquid crystal displays rely on rod-shaped polar molecules. Their dipoles let an electric field twist their orientation, which changes how they transmit polarised light.

Structure of solids

In crystals of polar molecules, the molecules pack so that opposite ends of neighbouring dipoles are close together, maximising the attraction.

Related forces

Dipole–dipole forces belong to a family of interactions involving dipoles:

• Ion–dipole: between an ion and a polar molecule — stronger, and responsible for salts dissolving in water (see ion–dipole forces).
• Dipole–induced dipole: a polar molecule distorts the electron cloud of a non-polar neighbour. This is why small amounts of non-polar gases such as O₂ dissolve in water.
• Hydrogen bonds: a particularly strong dipole–dipole attraction involving H bonded to N, O or F.

Writing a full exam answer

Question: Propanone (CH₃COCH₃) and butane (C₄H₁₀) have almost the same molar mass. Propanone boils at 56 °C and butane at −1 °C. Explain the difference.

Model answer: Both substances are made of simple molecules, so boiling only overcomes the forces between molecules, not the covalent bonds within them. The two molecules have a similar number of electrons (32 and 34), so the London forces between their molecules are of similar strength. However, propanone contains a polar C=O bond, and because of the molecule’s shape the bond dipole is not cancelled, so propanone molecules have a permanent dipole. Propanone therefore has permanent dipole–dipole forces between its molecules in addition to London forces. Butane is almost non-polar and has London forces only. More energy is needed to overcome the combined intermolecular forces in propanone, so it has the higher boiling point.

Notice the four moves in the answer: state that only intermolecular forces are overcome, compare London forces (similar size), identify the extra dipole–dipole forces and where they come from, and link the stronger total attraction to the energy needed and the boiling point. The same structure works for almost any comparison where the molecules are similar in size but differ in polarity.

Common misconceptions

• “Polar molecules don’t have London forces.” Every molecule has London forces.
• “The most polar molecule always has the highest boiling point.” Size can outweigh polarity (HCl → HI).
• “Dipole–dipole forces are bonds.” They’re attractions between molecules, far weaker than bonds.
• “CO₂ has dipole–dipole forces because C=O is polar.” CO₂ is linear and non-polar overall, so it doesn’t.

Key takeaways

• Permanent dipole–dipole forces are attractions between the δ+ and δ− ends of polar molecules.
• They’re typically about 5–25 kJ mol⁻¹ — stronger than London forces for small molecules, weaker than hydrogen bonds.
• Polar molecules have London forces too; dipole–dipole forces are extra.
• They explain higher boiling points of polar molecules of similar size (propane vs ethanal, cis vs trans).
• When sizes differ a lot, London forces may dominate.

Practise with intermolecular forces practice questions.