Van der Waals Forces: What the Term Really Covers

“Van der Waals forces” is one of the most confusingly used terms in school chemistry. Some textbooks use it to mean only London dispersion forces. Others use it for all intermolecular forces between neutral molecules. Still others include hydrogen bonds; some specifically exclude them. This article sorts out what the term covers, where it comes from, and how to use it safely in an exam.

Who was van der Waals?

Johannes Diderik van der Waals was a Dutch physicist. In his 1873 doctoral thesis, he tried to explain why real gases don’t behave exactly like the “ideal gas” of simple theory. He proposed two corrections: gas molecules take up space of their own, and they attract each other. The attractive forces he inferred — before anyone understood their origin — came to bear his name. He received the Nobel Prize in Physics in 1910.

The three components

In the modern, broad sense (the one used by IUPAC and in most physical chemistry), van der Waals forces are the attractive (and repulsive) forces between neutral molecules that don’t come from covalent bonding or from ions. The attractive part is usually divided into three types, each named after the scientist who explained it:

Component	Name	Between	Origin
Permanent dipole – permanent dipole	Keesom force	Two polar molecules	Attraction between fixed dipoles
Permanent dipole – induced dipole	Debye force	A polar and any molecule	A polar molecule distorts a neighbour’s electron cloud
Instantaneous dipole – induced dipole	London (dispersion) force	Any two molecules	Temporary electron fluctuations induce dipoles

Keesom forces (dipole–dipole)

Polar molecules such as HCl and propanone have permanent dipoles, and they attract when δ+ ends sit near δ− ends. In a liquid, molecules are tumbling, so this alignment is only partial; the effect weakens as temperature rises (see dipole–dipole forces).

Debye forces (dipole–induced dipole)

A polar molecule’s charged ends can distort the electron cloud of a neighbouring molecule — even a non-polar one — creating an induced dipole that’s attracted back. This explains why non-polar gases such as oxygen and nitrogen dissolve slightly in water: water’s dipoles induce small dipoles in O₂ and N₂. It’s this small solubility of oxygen — around 9 mg per litre at 20 °C — that lets fish breathe.

London forces (dispersion)

Moving electrons create instantaneous dipoles in all atoms and molecules, which induce dipoles in their neighbours. These are the only attractive forces between non-polar molecules and noble gas atoms, and they grow with the number of electrons and the surface area of contact (see London dispersion forces).

For most molecules, London forces make the largest contribution to the total van der Waals attraction — even for polar molecules like HCl and HBr. Only for small, highly polar molecules such as water does the dipole part dominate.

Where do hydrogen bonds fit?

Hydrogen bonds are usually treated as a separate, stronger category of intermolecular force, not as van der Waals forces (see hydrogen bonding). They involve a particularly strong interaction between a δ+ hydrogen bonded to N, O or F and a lone pair, with some partial covalent character and a strong preference for a straight-line geometry. IUPAC’s definition of van der Waals forces explicitly excludes hydrogen bonds and bonds involving ions.

Why the confusion in textbooks?

Many school syllabuses simplify the picture:

• Some use “van der Waals forces” to mean London forces only, and use “permanent dipole–dipole” and “hydrogen bonding” as separate categories. Under this usage, a statement like “iodine has van der Waals forces between its molecules” is fine, but “HCl has van der Waals forces and dipole–dipole forces” treats them as different things.
• Others use it for all three (Keesom + Debye + London).

Safe exam advice: use the specific name — “London (dispersion) forces”, “permanent dipole–dipole forces” or “hydrogen bonds” — rather than “van der Waals forces”. Examiners always accept the specific terms, and you avoid any ambiguity. If your syllabus defines the term, follow its definition.

The repulsive part and van der Waals radii

Van der Waals interactions aren’t only attractive. When two molecules get very close, their electron clouds start to overlap, and strong repulsion sets in. The balance between attraction (dominant at moderate distances) and repulsion (dominant at short distances) sets how close molecules approach each other in liquids and solids.

This gives each atom a van der Waals radius: half the distance between the nuclei of two identical atoms that are touching but not bonded. Van der Waals radii are always larger than covalent radii:

Element	Covalent radius (pm)	Van der Waals radius (pm)
H	31	120
C	76	170
N	71	155
O	66	152
Cl	102	175

Space-filling molecular models are built using van der Waals radii, showing the “surface” of a molecule that other molecules can’t penetrate.

Van der Waals and real gases

Van der Waals’ original goal was to improve the ideal gas equation, pV = nRT, which assumes gas particles have no volume and no attractions. His equation adds two corrections:

(p + a(n/V)²)(V − nb) = nRT

• a corrects for attractions between molecules, which reduce the pressure a real gas exerts on its container. Gases with stronger intermolecular forces have larger values of a.
• b corrects for the volume of the molecules themselves, which reduces the free space available.

Real gases behave most like ideal gases at high temperature and low pressure, where molecules are far apart and moving fast, so attractions and molecular volume matter least. At low temperatures and high pressures, van der Waals forces become important — and eventually cause the gas to condense into a liquid. You can explore ideal-gas calculations with the ideal gas law calculator.

Van der Waals forces in action

• Condensation of gases: all gases, even helium, liquefy at low enough temperatures because of London forces.
• Graphite: layers held together by van der Waals forces slide easily (see giant covalent structures).
• Gecko feet: millions of tiny contacts add up to enough van der Waals attraction to support the animal’s weight.
• Protein folding and drug binding: many small van der Waals contacts between closely fitting surfaces help hold molecules together (see how drugs are designed).
• Adhesives and tapes: some pressure-sensitive adhesives rely largely on van der Waals forces between the adhesive and the surface.

Common misconceptions

• “Van der Waals forces only exist in non-polar molecules.” In the broad sense, they act between all molecules.
• “Van der Waals forces and London forces are always the same thing.” Only in some textbook definitions.
• “Hydrogen bonds are a type of van der Waals force.” Usually classified separately.
• “Van der Waals forces are always tiny.” For large molecules and large contact areas, they can be substantial.

Key takeaways

• Van der Waals forces (broad definition) = Keesom (dipole–dipole) + Debye (dipole–induced dipole) + London (dispersion) forces between neutral molecules.
• Some syllabuses use the term for London forces only — when in doubt, use the specific force name.
• Hydrogen bonds and ion interactions are usually classed separately.
• Van der Waals radii describe how close non-bonded atoms can approach.
• The van der Waals equation corrects the ideal gas law for molecular attractions (a) and volume (b).