Change of State
Matter exists in distinct phases — solid, liquid, and gas — and can transition between them when energy is added or removed. These transitions, called changes of state or phase transitions, include melting (solid → liquid), freezing (liquid → solid), vaporisation (liquid → gas), condensation (gas → liquid), sublimation (solid → gas), and deposition (gas → solid). Understanding these processes is fundamental to thermal physics because they reveal how energy is stored and transferred at the molecular level. A crucial and often counterintuitive fact about phase changes is that temperature does not change while a substance is undergoing a phase transition — even though energy is continuously being added or removed. For example, a pot of boiling water stays at 100 °C until every last drop has evaporated, no matter how fiercely you heat it. The energy being supplied is not raising the temperature; instead, it is breaking the intermolecular bonds holding molecules together in the liquid phase. This hidden energy is called latent heat. Heating curves are graphs of temperature versus heat energy supplied (or time, at constant heating rate) that make these ideas visual. They show flat horizontal plateaus at melting and boiling points, separated by sloping sections where temperature rises. Reading a heating curve carefully lets you extract the specific heat capacity of each phase and the latent heats of fusion and vaporisation — quantities that are essential in engineering, climate science, cooking, and medicine.
- Heating a Solid — Kinetic Energy Increases
- Melting Plateau — Latent Heat of Fusion
- Heating the Liquid
- Boiling Plateau — Latent Heat of Vaporisation
- Heating the Gas
- Cooling and Reverse Transitions
Heating a Solid — Kinetic Energy Increases
When heat is added to a solid below its melting point, molecules vibrate more vigorously about their fixed lattice positions. The temperature rises according to Q = mcΔT, where c is the specific heat capacity of the solid phase. On a heating curve this appears as an upward slope.
Melting Plateau — Latent Heat of Fusion
At the melting point the temperature stops rising. All incoming energy goes into breaking intermolecular bonds, converting the ordered lattice into a mobile liquid. The energy required per unit mass is the specific latent heat of fusion, L_f. Only once all the solid has melted does the temperature begin rising again. The flat region on the heating curve at this point is called a plateau.
Heating the Liquid
With the solid fully melted, adding more heat raises the liquid's temperature again. Liquids generally have a different specific heat capacity than their solid form (water is unusual in that ice has a lower c than liquid water). The slope of the heating curve therefore changes at the end of the melting plateau.
Boiling Plateau — Latent Heat of Vaporisation
At the boiling point a second, typically longer, plateau appears. Vaporisation requires breaking nearly all remaining intermolecular attractions, so the specific latent heat of vaporisation, L_v, is always much larger than L_f for the same substance. For water, L_v ≈ 2260 kJ kg⁻¹ compared with L_f ≈ 334 kJ kg⁻¹ — almost seven times greater.
Heating the Gas
Once fully vaporised, adding heat raises the gas temperature. Gas molecules now move freely and rapidly; their specific heat capacity at constant pressure differs from both solid and liquid values. The heating curve slopes upward again, usually more steeply than the liquid section because gases typically have lower specific heat capacities per unit mass.
Cooling and Reverse Transitions
Every phase transition is reversible. Removing heat causes condensation and then freezing, and the same amounts of latent heat are released rather than absorbed. This is why steam burns are so much more severe than boiling-water burns at the same temperature — condensing steam releases an extra 2260 kJ per kilogram as it returns to liquid water on your skin.
The Key Equation
- Q — Heat energy (J): The total thermal energy transferred during the phase change. No temperature change occurs while this energy is being supplied or removed.
- m — Mass (kg): The mass of the substance undergoing the phase transition.
- L — Specific latent heat (J kg⁻¹): The energy required to change the phase of 1 kg of substance without changing its temperature. Use L_f for fusion (melting/freezing) and L_v for vaporisation (boiling/condensing).
Worked Example
Problem. A kettle contains 0.500 kg of water at exactly 100 °C. How much additional energy must be supplied to completely boil away all the water? (Specific latent heat of vaporisation of water: L_v = 2.26 × 10⁶ J kg⁻¹)
Answer. 1130 kJ (1.13 × 10⁶ J)
Origins · 1 minJoseph Black and how this was discovered
Joseph Black, a Scottish physicist and chemist, was the first to systematically study and quantify latent heat in the 1750s–1760s. He showed experimentally that ice absorbs a large, fixed quantity of heat when melting without any change in temperature, coining the term 'latent heat' (from Latin latere, 'to lie hidden'). He also discovered specific heat capacity, laying the quantitative foundation for calorimetry and all of thermal physics. Black's meticulous measurements were so influential that James Watt — his colleague at the University of Glasgow — credited Black's latent heat work as the key insight that led him to dramatically improve the efficiency of the steam engine, kickstarting the Industrial Revolution.
Related Concepts
- Specific Heat Capacity — Specific heat capacity governs the sloping sections of the heating curve between phase changes, and must be understood alongside latent heat to fully interpret heating curves.
- Thermal Energy Transfer — Phase changes involve large energy transfers; understanding conduction, convection, and radiation explains how that energy reaches or leaves a substance.
- Kinetic Theory of Matter — Latent heat makes sense at the molecular level through kinetic theory — phase transitions involve overcoming intermolecular forces, not speeding up molecules.