Introduction
Not all attractions between particles come from full ionic or covalent bonds. Molecules and atoms also experience weaker, short-range attractions called van der Waals forces. These forces explain why substances made of individual molecules (like O₂ or I₂) can still condense into liquids and solids, and why noble gases can be liquefied at all, even though their atoms don't form chemical bonds with each other. Van der Waals forces is actually an umbrella term covering three related types of intermolecular attraction (Talbot, p. 170).
The Three Types of Van der Waals Forces
1. London (dispersion) forces
These exist between all atoms and molecules, polar or non-polar. They arise because electrons are constantly moving, so at any instant a molecule may have a temporary, uneven distribution of charge — an instantaneous dipole. This can induce a dipole in a neighbouring molecule, leading to a weak, short-lived attraction.
- Strength increases with the number of electrons (and therefore molar mass), since more electrons give a larger, more polarizable electron cloud (Talbot, p. 112).
- Molecular shape matters too: long, thin molecules can line up side-by-side and have more contact area, giving stronger dispersion forces than compact, spherical molecules of similar mass.
2. Dipole–dipole forces
These occur between polar molecules, where the permanent δ+ end of one molecule attracts the permanent δ− end of another (Talbot, p. 170). These are generally stronger than dispersion forces between molecules of similar size, because the attraction is constant rather than momentary.
3. Dipole–induced dipole forces
A polar molecule can distort the electron cloud of a nearby non-polar molecule, creating a temporary induced dipole in it. The two then attract each other (Talbot, p. 170). A related idea is ion–induced dipole interaction, where an ion (e.g. Fe²⁺) polarizes a non-polar molecule such as O₂; the strength here depends on the ion's charge and how easily the neighbouring molecule is polarized (Talbot, p. 170).
Worked Example 1: Comparing boiling points of halogens
F₂, Cl₂, Br₂, and I₂ are all non-polar diatomic molecules, so only London forces act between them. Going down the group, the number of electrons increases, so London forces get stronger. This correctly predicts the trend in boiling points: F₂ (gas) < Cl₂ (gas) < Br₂ (liquid) < I₂ (solid).
Worked Example 2: Why does Cl₂ deviate more from ideal gas behaviour than NH₃?
Although Cl₂ is non-polar and NH₃ is polar (with hydrogen bonding), Cl₂ shows a greater deviation from ideal gas behaviour. This is because Cl₂ has more electrons, so its collective London forces are actually stronger than the hydrogen bonding in NH₃ (Talbot, p. 112). This shows London forces shouldn't be dismissed as always "weak" — with enough electrons, they can outcompete other intermolecular forces.
Worked Example 3: Van der Waals radius
For noble gases, which don't form covalent bonds, atomic size is measured using the van der Waals radius — half the distance between nuclei of two neighbouring, non-bonded atoms (e.g. in solid argon) (Talbot, p. 243). This is different from covalent or metallic radii, which rely on actual bonding.
Key Takeaways
- Van der Waals forces is a collective term for London (dispersion), dipole–dipole, and dipole–induced dipole forces.
- London forces exist in all particles and increase with more electrons (molar mass) and with molecule shape/surface area.
- Dipole–dipole forces only occur between polar molecules and are usually stronger than London forces for similarly-sized molecules.
- Strong London forces (e.g. in Cl₂) can sometimes outweigh weaker types of intermolecular attraction, affecting gas behaviour and boiling points.
- The van der Waals radius describes atomic size for non-bonded atoms, such as noble gases.