Vapour Pressure and Boiling: Concepts Explained Simply

Even in a closed bottle of water, some molecules at the surface escape into the air as vapour while vapour molecules rejoin the liquid. When these two rates balance, the vapour pushes on the walls with a steady pressure — the vapour pressure. It depends only on the liquid and its temperature, not on how much liquid is present.

This matters because boiling is defined through vapour pressure: a liquid boils when its vapour pressure equals the surrounding air pressure. That is why water boils at 100 °C at sea level but at about 71 °C on Mount Everest — the air pressure up there is much lower, so the vapour pressure catches up sooner.

Key Concepts

• Vapour pressure: the pressure exerted by a vapour in equilibrium with its liquid at a given temperature, in mmHg, torr or kPa. Higher temperature means higher vapour pressure.
• Intermolecular forces: strong attractions (hydrogen bonding in water) hold molecules back, giving low vapour pressure; weak attractions (diethyl ether) give high vapour pressure and fast evaporation.
• Boiling point: the temperature at which vapour pressure equals external pressure, so bubbles of vapour can form throughout the liquid.
• Normal boiling point: the boiling point at exactly 1 atm (101.3 kPa) external pressure — 100 °C for water.
• Evaporation vs boiling: evaporation happens at the surface at any temperature; boiling happens throughout the liquid but only at the boiling point.
• Clausius–Clapeyron equation: ln(P₂/P₁) = −ΔHvap/R × (1/T₂ − 1/T₁), relating vapour pressure to temperature (ΔHvap = molar enthalpy of vaporisation).

Formulas

• Clausius–Clapeyron: ln(P₂/P₁) = −(ΔHvap/R)(1/T₂ − 1/T₁), with T in kelvin, ΔHvap in J/mol, R = 8.314 J/mol·K.
• Boiling condition: Pvapour(Tboil) = Pexternal.

The equation assumes ΔHvap is constant over the temperature range, which is approximately true for small ranges.

Worked Example

On Mount Everest the air pressure is about 0.33 atm. At what temperature does water boil there? Use ΔHvap = 40.7 kJ/mol and the normal boiling point 373 K.

1. Set P₁ = 1.00 atm at T₁ = 373 K, P₂ = 0.33 atm at unknown T₂.
2. ln(P₂/P₁) = ln(0.33) = −1.1087.
3. −1.1087 = −(40700/8.314)(1/T₂ − 1/373) = −4895.4 × (1/T₂ − 0.002681).
4. Divide: 1/T₂ − 0.002681 = 1.1087/4895.4 = 0.0002265.
5. 1/T₂ = 0.0029075, so T₂ = 343.9 K = 70.8 °C.

Water boils at about 71 °C on Everest — which is why cooking takes so long there.

Common Misconceptions

• The boiling point of a liquid is a fixed property. It shifts with external pressure: lower pressure means a lower boiling point, which is also why pressure cookers (higher pressure) cook food faster.
• Bubbles in boiling water are made of air. The bubbles are pure water vapour forming inside the liquid — that is precisely what boiling means.
• High vapour pressure means strong intermolecular forces. It means the opposite: weakly attracted molecules escape easily, so volatile liquids like ether have high vapour pressure.

Real-World Applications

• Pressure cookers raise the boiling point to cook food faster at high altitudes.
• Vacuum distillation boils heat-sensitive liquids (like essential oils) at low temperatures.
• Refrigerators cycle refrigerants chosen for the right vapour-pressure curve.
• Perfume design balances volatile top notes against long-lasting base notes via vapour pressure.
• Weather and humidity: evaporation rates that drive the water cycle depend on vapour pressure.

Vapour pressure rises with temperature because more molecules can escape the liquid's attractions; boiling begins exactly when it matches the outside pressure. Everything from mountain cooking to perfume to refrigeration is this one idea at work.