van der Waals forces
/ van der VAHLS /
A gecko walks straight up a pane of glass, hanging on by nothing but the faint stickiness between the molecules of its toe pads and the molecules of the glass. That whisper-soft, ever-present attraction between neutral molecules is what we call van der Waals forces — the same gentle clinginess that lets dust cling to a screen and that pulls real gases slightly together.
Van der Waals forces are the collective name for the weak attractions between molecules that do not depend on full electrical charges or chemical bonds. They come in three flavors: London dispersion forces (from fleeting, momentary dipoles in any molecule), dipole–dipole forces (between molecules with a permanent lopsided charge), and dipole-induced-dipole forces (a permanent dipole nudging a neighbor). Strictly, chemists usually count hydrogen bonding as a separate, stronger category, though it is sometimes lumped in.
They matter because they are universal: every pair of molecules, however bland, feels at least the dispersion part of the van der Waals attraction. This is the same correction Johannes van der Waals added to the ideal-gas law to explain why real gases liquefy. The honest caveat is that any single van der Waals contact is feeble; their importance comes from acting over huge numbers of atoms at once, as on the gecko's foot.
Two argon atoms, with no charge and no chemical bond between them, still attract weakly enough that argon liquefies when cooled to minus 186 °C — pure van der Waals force at work.
Even noble-gas atoms attract through van der Waals forces.
Usage varies: some textbooks use 'van der Waals forces' for all intermolecular forces except hydrogen bonding, others restrict it to dispersion plus dipole forces, and a few use it loosely for dispersion alone. When precision matters, name the specific component you mean.