Vapor pressure and boiling
When a liquid is placed in a closed container, some surface molecules escape as vapor until evaporation and condensation balance in dynamic equilibrium. The pressure the vapor then exerts is the vapor pressure. As temperature rises, molecules gain kinetic energy and vapor pressure climbs steeply. Boiling begins when the vapor pressure equals the external pressure, so bubbles can form throughout the liquid; that temperature is the boiling point. The boiling point at 1 atm (101.325 kPa) is the normal boiling point, 100 °C for water.
How external pressure shifts the boiling point
- Lower pressure: atmospheric pressure falls with altitude, so liquids boil at a lower vapor pressure and thus a lower temperature. On the summit of Everest (about 0.33 atm) water boils at about 70 °C, and food cooks poorly.
- Higher pressure: a sealed pressure cooker raises internal pressure to about 2 atm, so water boils at about 120 °C and food cooks faster. Hospital autoclaves use the same principle.
- Intermolecular forces: at the same pressure, liquids with strong hydrogen bonding such as water have low vapor pressure and high boiling points, while ethanol or ether boil lower.
The Clausius–Clapeyron equation
Clapeyron (1834), developing Carnot's theory of heat engines, gave an equation for the slope of a phase-transition curve, which Clausius later grounded in thermodynamics. Treating the vapor as an ideal gas and neglecting the liquid's volume gives the approximation:
ln(P₂/P₁) = −(ΔHvap/R)(1/T₂ − 1/T₁)
Here ΔHvap is the molar enthalpy of vaporization and R the gas constant. Knowing the vapor pressure at one temperature lets you estimate the boiling point at another pressure.
Reading a phase diagram
On a pressure–temperature phase diagram, the vapor-pressure curve separating liquid and gas runs from the triple point (0.01 °C and 611 Pa for water) to the critical point (about 374 °C and 22.1 MPa). Every point on the curve is the boiling point at that pressure. Beyond the critical point, liquid and gas merge into a supercritical fluid.
Applications
- Vacuum distillation: lowering pressure lowers boiling points, so heat-sensitive substances can be separated without decomposing, as in oil refining and rotary evaporators.
- Freeze-drying and low-temperature concentration: removing water at low pressure and temperature preserves foods and medicines.
- Steam power and refrigeration: boilers and refrigerant cycles move heat by exploiting the pressure dependence of boiling points.
References
- Clapeyron, É. (1834). Mémoire sur la puissance motrice de la chaleur. Journal de l'École Royale Polytechnique, 14(23), 153–190.
- Atkins, P., de Paula, J., & Keeler, J. (2018). Atkins' Physical Chemistry (11th ed.). Oxford: Oxford University Press.