intermolecular force
/ IN-ter-mo-LEK-yoo-lar /
Pour water into a glass and it stays put as a puddle; let go of a balloon full of air and the gas rushes out and scatters everywhere. The difference is that water molecules cling to one another while air molecules barely do. That gentle pull between separate molecules is the intermolecular force — the reason matter can be liquid or solid at all instead of always flying apart as gas.
An intermolecular force is the attraction (and, at very short range, repulsion) acting between whole molecules, as opposed to the strong chemical bonds that hold the atoms inside a single molecule together. These forces are much weaker than chemical bonds — typically a few percent of a bond's strength — but they are what decide whether a substance boils at minus two hundred degrees or stays liquid at room temperature. They include dispersion forces, dipole–dipole attraction, hydrogen bonding, and ion–dipole interactions.
Why they matter: nearly every everyday property of liquids and solids — boiling point, melting point, surface tension, viscosity, solubility, the shape a protein folds into — traces back to how strongly molecules attract their neighbors. The honest caveat is that the word 'force' here covers a family of different mechanisms of different strengths, so naming the dominant one for a given substance is half the work.
Water boils at 100 °C while methane boils at minus 162 °C, even though their molecules weigh about the same — the water molecules attract one another far more strongly through intermolecular forces.
Stronger intermolecular forces mean a higher boiling point.
Do not confuse intermolecular forces (between molecules) with intramolecular bonds (within a molecule). When water boils, you break intermolecular forces — the molecules separate but each H₂O stays intact; breaking the actual O–H bonds takes far more energy.