What is a phase change?
A phase is a state of matter: solid, liquid or gas. A phase change is a move from one state to another. Nothing new is made. Water stays water.
- Melting (fusion): solid → liquid. Freezing: liquid → solid.
- Vaporisation: liquid → gas (by evaporation or boiling). Condensation: gas → liquid.
- Sublimation: solid → gas directly (dry ice, naphthalene balls). Deposition: gas → solid (frost).
Changes that need heat (melting, vaporisation, sublimation) are endothermic. Changes that give out heat (freezing, condensation, deposition) are exothermic.
Latent heat: why the temperature waits
When ice melts, you keep heating but the thermometer stays at 0 °C. Where does the heat go? It pulls particles apart against their attraction. It does not make them move faster, so the temperature does not rise. This hidden heat is called latent heat (latent means hidden).
Q = m × L
- Specific latent heat of fusion of water: Lf ≈ 334 kJ/kg (3.34 × 105 J/kg).
- Specific latent heat of vaporisation of water: Lv ≈ 2260 kJ/kg (2.26 × 106 J/kg).
Lv is much bigger than Lf. To make a gas, particles must be pulled fully apart, not just loosened. That is why steam at 100 °C burns worse than water at 100 °C: steam gives out its large latent heat when it condenses on skin.
Heating curve
A graph of temperature against heat added has rising parts (one phase warming up) and flat parts (a phase change). The flat part for boiling is longer than for melting because Lv > Lf.
Vapour pressure, volatility and boiling
In a closed bottle, some fast particles leave the liquid surface and become vapour. Some vapour particles fall back. Soon both rates are equal (a dynamic equilibrium). The pressure of the vapour now is the vapour pressure.
- Hotter liquid → more fast particles → higher vapour pressure.
- Weaker attraction between particles → easier escape → higher vapour pressure. Such liquids are volatile (petrol, acetone, ethanol).
Boiling happens when the vapour pressure equals the outside pressure. Then bubbles of vapour can form inside the liquid and not be squashed. The boiling point at 1 atm (101.3 kPa) is the normal boiling point: 100 °C for water, about 78 °C for ethanol, about 56 °C for acetone.
Evaporation vs boiling
Evaporation happens at any temperature, only at the surface, and slowly. Boiling happens at one fixed temperature (for a given pressure), all through the liquid, with bubbles.
Pressure changes the boiling point
Lower outside pressure → boiling point falls (mountains). Higher pressure → boiling point rises (pressure cooker, about 120 °C at 2 atm).
Try it at home
Put a thermometer in a pan of water with some ice and heat it slowly. Write the temperature every minute. You will see it stay near 0 °C while ice is left, then rise, then stay near 100 °C while it boils. Plot your own heating curve.
Phase diagrams, triple point and critical point
A phase diagram shows which phase is stable at each temperature (x-axis) and pressure (y-axis).
- Three areas: solid, liquid, gas.
- Three lines: melting line (solid–liquid), vapour pressure line (liquid–gas), sublimation line (solid–gas). On a line, two phases are in equilibrium.
- Triple point: where the three lines meet. All three phases exist together. For water: 0.01 °C and 0.006 atm (611 Pa).
- Critical point: the end of the liquid–gas line. Above it, liquid and gas cannot be told apart: a supercritical fluid. For water: 374 °C and 218 atm. For CO2: 31 °C and 73 atm (supercritical CO2 is used to remove caffeine from coffee).
Water is unusual
For most substances the melting line leans to the right. For water it leans slightly to the left, because ice is less dense than water. Squeezing ice can melt it a little.
CO2: why dry ice sublimes
The triple point of CO2 is at 5.1 atm. At normal pressure (1 atm) it is below the triple point, so solid CO2 turns straight into gas at −78 °C.
Key formulas and definitions
- Q = m × L (heat for a phase change, no temperature change)
- Q = m × c × ΔT (heat to warm one phase)
- Water: L_f ≈ 3.34 × 10^5 J/kg, L_v ≈ 2.26 × 10^6 J/kg, c ≈ 4200 J/(kg·K)
- Boiling: vapour pressure = outside pressure
- Water: triple point 0.01 °C, 611 Pa; critical point 374 °C, 22.1 MPa
Worked examples
1. How much heat melts 2 kg of ice at 0 °C? (L_f = 3.34 × 10^5 J/kg)
Q = m × L_f = 2 × 3.34 × 10^5 = 6.68 × 10^5 J = 668 kJ. The temperature stays at 0 °C the whole time.
2. How much heat turns 0.5 kg of water at 100 °C into steam at 100 °C? (L_v = 2.26 × 10^6 J/kg)
Q = 0.5 × 2.26 × 10^6 = 1.13 × 10^6 J = 1130 kJ.
3. How much heat changes 1 kg of ice at 0 °C into water at 20 °C? (c = 4200 J/kg K)
Step 1, melt: Q1 = 1 × 3.34 × 10^5 = 334 000 J. Step 2, warm: Q2 = 1 × 4200 × 20 = 84 000 J. Total = 418 000 J = 418 kJ.
4. Find the total heat to turn 0.2 kg of ice at −10 °C into steam at 100 °C. (c_ice = 2100, c_water = 4200 J/kg K)
Warm ice: 0.2 × 2100 × 10 = 4 200 J. Melt: 0.2 × 3.34 × 10^5 = 66 800 J. Warm water: 0.2 × 4200 × 100 = 84 000 J. Boil: 0.2 × 2.26 × 10^6 = 452 000 J. Total = 607 000 J ≈ 607 kJ. Most of it is for boiling.
5. At a hill station the air pressure is 0.7 atm. Will water boil above or below 100 °C? Why?
Below 100 °C (about 90 °C). Boiling needs vapour pressure = outside pressure. With less outside pressure, a lower temperature is enough.
6. On a phase diagram, a sample of water is at 0.01 °C and 611 Pa. What do you see?
This is the triple point. Ice, liquid water and water vapour all exist together in equilibrium.
Common mistakes
- Thinking temperature rises while ice melts. It stays at 0 °C until all the ice has melted.
- Saying boiling happens only at 100 °C. Water boils at 100 °C only at 1 atm; the boiling point changes with pressure.
- Mixing up evaporation and boiling. Evaporation is surface-only and happens at any temperature.
- Using Q = m c ΔT during a phase change. During the change ΔT = 0, so use Q = m L.