Intermolecular Forces: Attractions Between Molecules

Between molecules, not within them

The word 'intermolecular' does heavy lifting: these forces act between molecules, not within them. Covalent and ionic bonds are intramolecular — they hold atoms together inside a particle. Intermolecular forces are the weaker attractions that make those particles cling to each other.

Boiling water makes the distinction physical. At 100 °C, water molecules gain enough energy to break free of each other and fly off as steam — but each departing molecule is still intact H2O. The hydrogen bonds between molecules broke; the O−H covalent bonds within each molecule did not. Producing hydrogen and oxygen gas would require electrolysis, an entirely different energy regime.

The energy scales tell the story numerically: intermolecular forces run roughly 2–30 kJ/mol, while covalent bonds demand 150–1000 kJ/mol to break. Confusing the two scales is the single most common error in this topic — keep 'between' and 'within' separated and the rest follows.

Meet the three van der Waals forces

Three forces share the van der Waals name, plus one honorary member:

• London dispersion forces — present between all molecules. Momentary uneven electron distributions create temporary dipoles that induce matching dipoles in neighbors.
• Dipole-dipole forces — between polar molecules only. The δ+ end of one molecule attracts the δ− end of another.
• Hydrogen bonding — a special, strong dipole attraction when H bonded to N, O, or F meets a lone pair on N, O, or F.
• Ion-dipole forces — between an ion and a polar molecule, as when water solvates Na+. Strongest of the set.

Each force gets its own deep-dive article after this overview. For now, learn the roster and the membership rules: dispersion is universal, dipole-dipole needs polarity, hydrogen bonding needs the N/O/F pattern, and ion-dipole needs actual ions.

Ranking strengths

As a starting guide, strength runs: ion-dipole > hydrogen bonding > dipole-dipole > London dispersion. Typical energies: hydrogen bonds 5–30 kJ/mol, dipole-dipole interactions a few kJ/mol, dispersion from under 1 kJ/mol in helium to tens of kJ/mol in large molecules.

But the ranking comes with a large asterisk: London dispersion grows with molecular size and can overtake everything below ion-dipole. Iodine (I2) is a solid at room temperature held together by dispersion alone, while hydrogen fluoride — a hydrogen bonder — is a gas. Size rewrites the leaderboard.

Use the ranking to compare similar-sized molecules, where it works beautifully: among the hydrogen halides' lighter members or among small organics, stronger force types mean higher boiling points. When sizes differ wildly, reach for polarizability instead of the ranking. The ranking is a default assumption, and defaults exist to be overridden by evidence.

Worked comparison: water vs hydrogen sulfide

One comparison proves that intermolecular forces — not molecular mass — control boiling points. Water: molar mass 18 g/mol, boiling point 100 °C. Hydrogen sulfide: molar mass 34 g/mol, boiling point −60 °C.

If mass decided boiling points, the heavier H2S would boil higher. It boils 160 degrees lower instead. The explanation is force type: water's O−H groups form strong hydrogen bonds, while H2S — sulfur is too large and diffuse for effective hydrogen bonding — manages only dipole-dipole attractions.

Walk the logic explicitly: heavier molecule + weaker forces → lower boiling point, therefore forces dominate mass here. This is the worked reasoning examiners expect: identify the force each substance can exert, rank the forces, and let the ranking predict the order. Mass is a tiebreaker for similar forces, never the main event.

'Like dissolves like'

Solubility obeys the slogan 'like dissolves like,' and intermolecular forces are the reason. Dissolving means replacing solvent-solvent and solute-solute attractions with solvent-solute ones — the swap only pays if the new attractions rival the old.

Polar water welcomes polar ethanol and ionic sodium chloride: hydrogen bonds and ion-dipole forces replace water's own network at comparable strength. Nonpolar oil finds no such bargain — water would have to surrender hydrogen bonds for weak dispersion contacts, so oil stays separate and floats.

Soap bridges both worlds. Each soap molecule carries a polar, ionic head that dives into water and a long nonpolar tail that dives into grease. The tails pry grease loose while the heads keep the whole assembly suspended — micelles of grease wrapped in water-compatible skins, rinsed away together. One molecule, two force preferences, and the basis of all cleaning.

The properties IMFs control

Beyond boiling points and solubility, intermolecular forces tune a family of everyday properties. Viscosity: stronger attractions make liquids thicker — honey's extensive hydrogen bonding makes it pour slowly, while weakly attracted gasoline flows freely.

Surface tension: water beads on wax because surface molecules, pulled inward by hydrogen bonds with no neighbors above, minimize their surface area. Vapor pressure: weakly attracted liquids evaporate readily — rubbing alcohol (mostly dispersion forces) dries fast, while hydrogen-bonded water lingers.

Adhesion and capillarity complete the set: water climbs a paper towel because adhesive forces to cellulose rival its own cohesion. Each property is a macroscopic readout of invisible attractions.

This overview mapped the territory. The three detail articles ahead — hydrogen bonding, dipole-dipole, and London dispersion — examine each force's mechanism, strength, and signature examples up close.

Common mistakes

Real-world applications

• Designing distillation separations by boiling point

• Choosing refrigerants with the right vapor pressure

• Formulating paints, inks, and adhesives

• Explaining why oil spills float and resist mixing