2.2.3 Melting, Boiling and Evaporation
Learning Objective
Describe how temperature, surface area and air movement over a surface affect evaporation, and explain the related phase‑change concepts required by the Cambridge IGCSE 0625 syllabus.
Key Definitions (Core Syllabus)
- Melting: The change of state from solid to liquid when the required latent heat of fusion is supplied. The temperature of the substance remains constant at its melting point until the whole solid has melted.
- Boiling: The rapid change of state from liquid to vapour throughout the bulk of the liquid when the required latent heat of vapourisation is supplied and the temperature reaches the boiling point. The temperature stays constant at the boiling point until all the liquid has vapourised.
- Evaporation: The escape of high‑energy molecules from the surface of a liquid at any temperature below the boiling point. The escaping molecules take away kinetic energy, so the remaining liquid cools.
- Condensation: The reverse of evaporation – vapour molecules lose kinetic energy, come together and form a liquid.
- Solidification (Freezing): The reverse of melting – particles lose kinetic energy, arrange into a regular, closely‑packed lattice and release the latent heat of fusion.
Reference Temperatures for Water (Standard Pressure)
- Melting point: 0 °C
- Boiling point: 100 °C
- These values are the benchmark figures used in many IGCSE exam questions.
Differences between boiling (bulk evaporation) and evaporation (surface phenomenon)
| Aspect | Boiling | Evaporation |
|---|
| Location of change | Occurs throughout the liquid (bulk) | Occurs only at the liquid–air interface (surface) |
| Temperature requirement | Must reach the boiling point of the liquid | Can occur at any temperature below the boiling point |
| Heat supplied | Latent heat of vapourisation + sensible heat to reach the boiling point | Only the latent heat needed for the molecules that escape |
| Visible signs | Rapid formation of bubbles throughout the liquid | Gradual loss of liquid; no bubbles unless the liquid is near its boiling point |
How Evaporation Cools
- High‑energy molecules leave the surface, taking their kinetic energy with them.
- The average kinetic energy of the remaining molecules falls, so the temperature of the liquid drops – the familiar “evaporative cooling” (e.g., sweating).
- The energy removed is the latent heat of vapourisation of the escaped molecules.
Factors Influencing the Rate of Evaporation
Quantitative (qualitative) description
The evaporation rate, E, can be expressed as
\$\$
E = k \, A \, \bigl(P{\text{sat}}(T) - P{\text{air}}\bigr) \, f(v)
\$\$
- k – proportionality constant (depends on the liquid’s nature)
- A – exposed surface area
- Psat(T) – saturation vapour pressure at temperature T (rises sharply with T)
- Pair – partial pressure of that vapour already present in the surrounding air
- f(v) – function representing the effect of air velocity v (larger v → larger f(v))
1. Temperature
- Raising the temperature increases Psat(T) exponentially (Clausius‑Clapeyron relation).
- More molecules have kinetic energy greater than the escape energy, so more molecules leave per second.
- Example: A glass of water on a sunny windowsill evaporates faster than the same glass in a shaded cupboard.
2. Surface Area
- Evaporation occurs only at the liquid–air interface.
- Doubling the exposed area roughly doubles the number of molecules that can escape per unit time, all else being equal.
- Example: A thin film of spilled milk on a tray dries quicker than the same volume collected in a deep bowl.
3. Air Movement (Wind)
- In still air the vapour layer above the surface quickly becomes saturated, reducing the concentration gradient that drives evaporation.
- Moving air continually replaces saturated air with drier air, keeping the gradient large and the rate high.
- Example: Clothes dry faster on a breezy day than on a calm day.
4. Ambient Pressure (optional)
- Lower external pressure reduces the vapour pressure that must be reached for molecules to escape.
- This is why water boils at a lower temperature on a mountain top.
Exam‑style Summary of the Three Core Factors
When a question asks how a factor influences evaporation, use the following concise structure:
- Temperature ↑ → kinetic energy ↑ → more molecules exceed the escape energy → evaporation rate ↑.
- Surface area ↑ → more molecules are directly exposed → number of escaping molecules per second ↑.
- Air movement ↑ → saturated layer removed → concentration gradient ↑ → evaporation rate ↑.
Summary Table
| Condition | Effect on Evaporation Rate | Reason |
|---|
| High temperature | Increase | Higher kinetic energy → more molecules escape |
| Low temperature | Decrease | Fewer molecules have sufficient energy |
| Large surface area | Increase | More molecules are at the interface |
| Small surface area | Decrease | Fewer molecules exposed |
| Strong air flow | Increase | Vapour removed, concentration gradient maintained |
| Still air | Decrease | Vapour builds up, reducing gradient |
| Low ambient pressure | Increase | Lower vapour pressure needed for escape |
Suggested Diagram
Include a single sketch that shows:
- (a) A shallow pan of water – large surface area.
- (b) A deep bowl – small surface area.
- (c) Airflow over a water surface (arrows indicating wind).
- (d) High‑energy molecules at the surface escaping, leaving the remaining liquid cooler.
Exam Tips
- Link each factor to kinetic energy, surface exposure, or concentration gradient when explaining its effect.
- Use the qualitative formula E = kA(Psat‑Pair)f(v) to justify quantitative‑style answers.
- Remember: Boiling = bulk evaporation at the boiling point; evaporation = surface phenomenon at any temperature.
- If a question mentions a temperature change of the liquid, mention the cooling effect of evaporation.
- For higher‑grade questions, note that the boiling point varies with ambient pressure (e.g., lower on a mountain).