interatomic forces
/ in-ter-uh-TOM-ik FOR-sez /
Imagine each atom is connected to its neighbours by tiny springs — pull them too far apart and the springs pull back; shove them too close and the springs push away. Atoms really do behave a bit like this, and the pushes and pulls between them are called interatomic forces. They are the glue that decides whether matter holds together at all.
These forces come from electricity at the atomic scale: the negatively charged electrons and positively charged nuclei of nearby atoms attract and repel one another. At a comfortable distance the net pull holds the atoms together, but squeeze them closer and the repulsion grows fiercely, which is why solids resist being crushed. The balance point — where pulling and pushing cancel — is the natural spacing the atoms prefer to sit at.
This matters because interatomic forces set almost every property of a material: how strong it is, how hot it must get to melt, how much it expands when warmed. A common misconception is that there is a single kind of bond; in fact there are several distinct flavours — ionic, covalent, metallic, and weaker ones like van der Waals — and which one rules decides whether you get diamond, salt, copper, or candle wax.
Diamond is the hardest natural material because its carbon atoms are bound by exceptionally strong, stiff forces, while the graphite in a pencil is made of the very same carbon — but with weak forces between its layers, so it flakes off onto paper at the lightest touch.
Same carbon atoms, very different forces: diamond is hard, graphite flakes.
Interatomic forces are both attractive and repulsive at once: attraction dominates at large spacing and repulsion at small spacing, so atoms settle where the two exactly balance rather than touching outright.