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Matter exists as solid (fixed shape), liquid (fixed volume), or gas (expands to fill container). Phase changes occur when temperature or pressure changes provide enough energy to overcome intermolecular forces. Melting (solid→liquid), boiling (liquid→gas), and sublimation (solid→gas) are endothermic; reverse processes are exothermic. Heat of fusion and vaporization quantify energy needed.
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 fundamentally a contest between these attractive forces pulling molecules together and kinetic energy (thermal motion) trying to fling them apart. In a solid, intermolecular forces win decisively — molecules vibrate in fixed positions within an ordered lattice. In a liquid, kinetic energy is high enough that molecules slide past each other but not high enough to escape the collective pull entirely. In a gas, kinetic energy overwhelms the attractions and molecules fly freely, filling whatever container they occupy.
A phase change happens at the tipping point where the balance shifts. When you heat ice, you add kinetic energy. At 0°C, the molecules have enough energy to break free of the rigid crystal lattice — this is melting. Crucially, temperature stays constant during a phase change even though you keep adding heat. That energy is not increasing molecular speed; it is being consumed entirely by breaking intermolecular attractions. The energy required to melt one mole of a substance is its heat of fusion (ΔH_fus). Continue heating the liquid water to 100°C, and molecules gain enough energy to escape the liquid surface entirely — boiling. The heat of vaporization (ΔH_vap) is always much larger than the heat of fusion because going from liquid to gas requires completely overcoming intermolecular forces, whereas melting only loosens the structure partially.
Sublimation — a solid converting directly to gas, as dry ice does — occurs when surface molecules gain enough energy to escape the lattice entirely without passing through the liquid phase. This happens when vapor pressure at the solid's surface exceeds atmospheric conditions that would otherwise stabilize a liquid. The reverse processes — freezing, condensation, and deposition — are exothermic because forming intermolecular attractions releases energy. Every phase change is thus a story told in the language of intermolecular forces: stronger forces mean higher melting and boiling points, larger heats of fusion and vaporization, and a greater reluctance to enter the gas phase.
The practical consequence is predictive power. If you know a substance has strong hydrogen bonding (like water), you can predict it will have an unusually high boiling point relative to its molecular weight. If a substance has only weak London dispersion forces (like methane), it will be a gas at room temperature. Phase diagrams, which you will encounter later, map out these relationships across all combinations of temperature and pressure, but the underlying logic is always the same: intermolecular forces versus kinetic energy.