Introduction
Not all bonds hold atoms together permanently inside molecules — some are much weaker forces that act between molecules. These are called van der Waals forces, and they explain why substances made of individual molecules can still be liquids or solids, why gases don't behave perfectly ideally, and why some molecules stick to each other more than others. Understanding these forces helps explain boiling points, solubility, and the physical states of many everyday substances.
What are van der Waals forces?
The term "van der Waals forces" is actually an umbrella term covering three types of intermolecular attraction:
- London (dispersion) forces — present between all molecules and atoms, even non-polar ones
- Dipole–dipole forces — between molecules with permanent dipoles
- Dipole–induced dipole forces — between a polar molecule and a non-polar one
(Talbot, p.170)
All three are much weaker than covalent, ionic, or metallic bonds, but they are essential for holding molecules together in liquids and solids.
London (dispersion) forces
These arise because electrons are constantly moving. At any instant, electron distribution around an atom or molecule may become uneven, creating a temporary/instantaneous dipole. This temporary dipole can induce a similar dipole in a neighbouring molecule, leading to a weak attraction.
Key point: London forces exist between all particles, including noble gas atoms and non-polar molecules like O₂ and Cl₂, because they only require electrons — and every atom has those.
Strength depends on:
- Number of electrons / molar mass — more electrons means a larger, more polarisable electron cloud, so stronger dispersion forces (Talbot, p.112)
- Shape of the molecule — longer, more elongated molecules have greater surface contact area, allowing stronger London forces than compact, spherical molecules of similar mass