States of Matter and Phase Transitions

Core Idea

Matter exists in three main states—solid, liquid, and gas—distinguished by how closely particles are packed and how freely they move. Phase transitions occur when energy addition or removal overcomes intermolecular forces. Melting, vaporization, sublimation, and their reverses are endothermic or exothermic processes.

Explainer

From your study of intermolecular forces, you know that molecules attract each other through dipole-dipole interactions, hydrogen bonds, and London dispersion forces. The state of matter a substance adopts is essentially a contest between these attractive forces pulling molecules together and the kinetic energy of the molecules trying to fly apart. In a solid, intermolecular forces win decisively — particles are locked into fixed positions, vibrating in place but unable to move past their neighbors. In a liquid, kinetic energy is high enough that particles slide past one another while remaining in close contact. In a gas, kinetic energy overwhelms the attractive forces entirely, and particles move independently with large spaces between them.

Phase transitions happen when the balance tips. When you heat a solid, you are adding kinetic energy to the particles. At the melting point, the added energy is just enough to overcome the forces holding particles in their fixed lattice positions, and the solid becomes a liquid. Crucially, during the transition itself, the temperature does not rise — all the energy being added goes into breaking intermolecular attractions rather than increasing particle speed. This is the enthalpy of fusion. The same principle applies at the boiling point, where the enthalpy of vaporization represents the energy needed to fully separate liquid-phase particles into the gas phase. Because vaporization requires overcoming all remaining intermolecular contact, it always demands more energy than melting.

The reverse processes release energy. Condensation (gas to liquid) and freezing (liquid to solid) are exothermic — the formation of intermolecular attractions releases the same energy that was required to break them. Sublimation is the direct transition from solid to gas, skipping the liquid phase entirely, and it requires energy equal to the sum of fusion and vaporization enthalpies. Dry ice (solid CO₂) sublimes at atmospheric pressure because CO₂'s weak London dispersion forces and low molecular symmetry make the liquid phase unstable under normal conditions.

The strength of a substance's intermolecular forces directly predicts its phase behavior. Water, with its extensive hydrogen bonding network, has an unusually high boiling point for its molecular weight. Methane, relying only on weak London forces, is a gas at room temperature. Comparing boiling points across a series of molecules is really comparing the strength of their intermolecular forces — a principle that connects this topic directly to everything you learned about molecular polarity and intermolecular attractions.