Hydrogen Bonds vs Van der Waals Forces: The Difference

Both hold molecules together from the outside, both are far weaker than a real bond — so why does one get a special name? Hydrogen bonds and van der Waals forces sit in the same family, but the gap in strength between them explains some of chemistry's biggest surprises, like why water boils at 100 °C when it "should" boil far colder.

The short answer: van der Waals forces are the weak, general attractions between molecules — mainly London dispersion forces and dipole–dipole forces. A hydrogen bond is a much stronger, special attraction that forms only when hydrogen is bonded to nitrogen, oxygen, or fluorine and is drawn to a lone pair on an N, O, or F in another molecule. Both are intermolecular forces; the hydrogen bond is just the strongest of the group.

Quick comparison at a glance

Feature	Van der Waals forces	Hydrogen bonds
What they include	London dispersion + dipole–dipole	A special, strong dipole–dipole case
Where they occur	Between all molecules	Only with H bonded to N, O, or F
Relative strength	Weak to moderate (~0.5–20 kJ/mol)	Stronger (~5–30 kJ/mol)
Depends on	Molecule size, polarity	Presence of N–H, O–H, or F–H
Everyday effect	Why noble gases can be liquefied; gecko grip	Water's high boiling point; DNA holding together

Both are still far weaker than the covalent or ionic bonds inside molecules — this whole comparison lives in the "between molecules" league.

A note on the naming (read this first)

Textbooks split two ways, so it's worth being clear. Van der Waals forces is an umbrella term for the weak intermolecular attractions — most courses use it to mean London dispersion forces and dipole–dipole forces. Hydrogen bonding is technically an especially strong kind of dipole–dipole attraction, but because it's so much stronger than the rest, most intro courses list it separately as its own category.

So when a question contrasts "hydrogen bonds vs van der Waals forces," it means: the strong special case vs the ordinary weaker attractions. That's the split we'll use here.

What van der Waals forces are

Under the umbrella sit two forces:

• London dispersion forces. Present in every molecule, even nonpolar ones. Electrons are always shuffling around, and for an instant a molecule can have a lopsided charge — a temporary dipole — that induces one in its neighbour. Weak and fleeting, but they add up: bigger molecules with more electrons have stronger dispersion forces (which is why iodine is a solid but chlorine is a gas).
• Dipole–dipole forces. Between polar molecules, the δ+ end of one is attracted to the δ− end of the next. Stronger than dispersion for molecules of similar size.

What a hydrogen bond is

A hydrogen bond needs a specific set-up. It forms when a hydrogen atom is covalently bonded to a small, very electronegative atom — N, O, or F — which strips so much electron density from the H that it becomes strongly δ+. That exposed hydrogen is then powerfully attracted to a lone pair on an N, O, or F of a neighbouring molecule.

Only N, O, and F work: they're electronegative and small enough to concentrate the charge. That's why "N, O, F" is worth memorising as the hydrogen-bond entry ticket.

Why this explains water

Water (H₂O) is small and light, so by size alone it "should" boil around −80 °C, like similar small molecules. Instead it boils at +100 °C. The reason is hydrogen bonding: every water molecule can hydrogen-bond to several neighbours through its O–H groups and oxygen lone pairs. Pulling those molecules apart to boil takes a lot of extra energy. The same force makes ice less dense than liquid water (so ice floats) and lets water climb up plant stems.

Worked examples

Identify the main intermolecular force:

• Two argon atoms (Ar): nonpolar → London dispersion only.
• HCl molecules: polar, no N/O/F–H → dipole–dipole (van der Waals).
• Water (H₂O): O–H present → hydrogen bonding (plus weaker dispersion underneath).
• Ammonia (NH₃): N–H present → hydrogen bonding.
• Methane (CH₄): nonpolar, H on carbon (not N/O/F) → dispersion only, no hydrogen bonds.

Common mistakes to avoid

• Thinking a hydrogen bond is a real chemical bond. It isn't — it's an intermolecular attraction between molecules, roughly a tenth the strength of the covalent bonds inside them.
• Expecting hydrogen bonds from any H. The hydrogen must be bonded to N, O, or F. C–H groups (as in methane) don't hydrogen-bond, so they rely on dispersion.
• Forgetting dispersion is everywhere. London dispersion forces act between all molecules, including ones that also hydrogen-bond — hydrogen bonds sit on top of them, not instead of them.

FAQ

What's the difference between a hydrogen bond and a van der Waals force?
Van der Waals forces are the general weak attractions (dispersion and dipole–dipole) between molecules. A hydrogen bond is a stronger, special case that only forms when H is bonded to N, O, or F.

Is a hydrogen bond a type of van der Waals force?
Technically it's a very strong dipole–dipole interaction, which falls under van der Waals in the strict sense. But it's so much stronger that most courses treat it as its own separate category.

Which is stronger, a hydrogen bond or a van der Waals force?
The hydrogen bond, typically several times stronger than ordinary dispersion or dipole–dipole forces — though all are much weaker than covalent or ionic bonds.

Why does water have such a high boiling point?
Hydrogen bonding. Each water molecule bonds to several neighbours, so it takes extra energy to separate them, pushing the boiling point far above what water's small size would predict.

The takeaway

Van der Waals forces are the everyday weak attractions between molecules (dispersion for all, dipole–dipole for polar ones); a hydrogen bond is the strong special case that appears only with H–N, H–O, or H–F. Both are intermolecular, both are weaker than real bonds — but the extra muscle of the hydrogen bond is exactly why water behaves the way it does.

Related → [Intramolecular vs Intermolecular Forces], [What Is a Polar Molecule?], and [Polar vs Nonpolar Bonds].