Phase Changes and Diagrams

Core Idea

Matter transitions between solid, liquid, and gas phases when energy is added or removed. Melting (solid→liquid), vaporization (liquid→gas), and sublimation (solid→gas) are endothermic; their reverses are exothermic. During a phase transition, temperature remains constant as added energy breaks intermolecular forces rather than increasing kinetic energy. Phase diagrams plot pressure vs temperature and show which phase is stable under given conditions. The triple point is the unique P-T condition where all three phases coexist in equilibrium. The critical point marks the end of the liquid-gas boundary — above it, a supercritical fluid exists with properties of both phases.

How It's Best Learned

Trace heating curves (temperature vs heat added) to see how temperature plateaus during phase transitions. Read phase diagrams by identifying regions, boundaries, and special points. Compare phase diagrams of water (negative solid-liquid slope) and CO₂ (positive slope) to understand how pressure affects melting.

Common Misconceptions

• Boiling point is not a fixed property of a substance — it depends on external pressure. Water boils below 100°C at high altitude because atmospheric pressure is lower.
• During a phase transition, added heat does not raise temperature. Students often expect temperature to rise continuously; the plateaus on heating curves reflect energy going into overcoming intermolecular forces.

Explainer

From thermochemistry, you know that enthalpy changes track heat flow, and from intermolecular forces, you know that molecules attract each other through dipole-dipole interactions, hydrogen bonds, and London dispersion forces. Phase changes are what happen when thermal energy either overwhelms these intermolecular attractions or loses the battle against them. Melting, boiling, and sublimation are endothermic because energy must be absorbed to pull molecules apart; freezing, condensation, and deposition are exothermic because energy is released as molecules settle into closer, more ordered arrangements.

The heating curve makes the energy story visible. When you heat ice from −20°C, the temperature rises steadily as the added energy increases molecular kinetic energy (the sloped portions). But at 0°C something striking happens: the temperature stops rising even though you're still adding heat. All the incoming energy goes into breaking the hydrogen bonds that hold the ice lattice together — this is the heat of fusion (ΔH_fus = 6.01 kJ/mol for water). Only after all the ice has melted does the temperature resume climbing. The same plateau occurs at 100°C during vaporization, except the heat of vaporization (ΔH_vap = 40.7 kJ/mol) is much larger because vaporization requires completely separating molecules from each other, not just disrupting a lattice. This is why it takes far more energy to boil water away than to melt ice — and why steam burns are so much worse than hot water burns, as the condensing steam releases all that stored energy onto your skin.

A phase diagram maps which phase is thermodynamically stable at each combination of pressure and temperature. The boundaries between regions are lines where two phases coexist in equilibrium — the solid-liquid boundary, the liquid-gas boundary, and the solid-gas boundary. The triple point is the unique temperature and pressure where all three phase boundaries meet and all three phases coexist simultaneously (for water: 0.01°C, 0.006 atm). The critical point marks the end of the liquid-gas boundary — above this temperature and pressure, the distinction between liquid and gas disappears, and a supercritical fluid exists with properties of both phases (supercritical CO₂ is used as a solvent in decaffeination).

Water's phase diagram has a famous anomaly: its solid-liquid boundary slopes to the left (negative slope), meaning that increasing pressure at constant temperature can melt ice. This happens because ice is less dense than liquid water — pressure favors the denser phase. For nearly every other substance, the solid-liquid line slopes to the right (positive slope) because the solid is denser. This quirk of water is why ice floats, why lakes freeze from the top down, and why ice skating works — pressure under the blade slightly lowers the melting point, though the effect is much smaller than commonly claimed. Reading a phase diagram is a matter of placing your finger at a P-T coordinate and seeing which region you're in, then tracing how phase changes occur as you move along a path of changing temperature or pressure.