Evaporation and boiling

In Leadville, Colorado, about 3,100 m up, air pressure is about 68 kPa and water boils at about 90 °C.

A puddle dries on a cool day without ever getting near 100 °C. A pot of water on the stove, on the other hand, sits quietly until it starts to bubble at about 100 °C. Both are vaporization, the change from liquid to gas, but they are not the same process. Evaporation happens only at the surface and at any temperature. Boiling happens all through the liquid, at one particular temperature that depends on the pressure above it.

Evaporation and vapour pressure

In any liquid the molecules have a range of speeds. A few at the surface are moving fast enough to break away from their neighbours and escape into the air. That is evaporation. Because the fastest molecules leave first, the ones left behind have a lower average energy, so the liquid cools. Sweat works this way. OpenStax estimates that evaporating 1.5 L of sweat in a hot climate removes about 3,600 kJ of heat from the body.

In a closed bottle, escaping molecules collect above the liquid, and some of them bump back into it. Eventually as many return as leave each second. The pressure of the vapour at that point is the vapour pressure of the liquid. It depends on the substance and the temperature, not on the size of the bottle. Liquids whose molecules attract weakly have high vapour pressures: at 20 °C methanol's is 11.9 kPa and ethanol's is 5.95 kPa. Heat a liquid and more molecules have enough energy to escape, so its vapour pressure rises.

Boiling depends on the air above

A liquid boils when its vapour pressure equals the pressure of the surroundings. Then bubbles of vapour can form inside the liquid without being crushed. The normal boiling point is the boiling point at 1 atm, which is 101.3 kPa. For water that is 100 °C. High in the mountains the air pressure is lower, so water boils at a lower temperature: in Leadville, where the pressure is about 68 kPa, it boils at about 90 °C. A pressure cooker does the opposite. It traps steam, raises the pressure and lets the water get hotter than 100 °C before it boils.

Vaporizing takes a lot of energy, more than melting, because every attraction to neighbouring molecules has to be broken, not only loosened. For water the enthalpy of vaporization is 44.01 kJ/mol at 25 °C and 40.67 kJ/mol at the boiling point. Compare that with 6.01 kJ/mol to melt ice: boiling a mole of water takes almost seven times as much energy as melting it. Condensation, gas to liquid, gives that energy back, which is why a scald from steam is worse than one from water at the same temperature.

Skipping the liquid

Some solids turn straight into gas. This is sublimation. Dry ice, solid carbon dioxide, slowly vanishes at room temperature without leaving a puddle. NIST data show why: carbon dioxide can only be liquid above 5.185 bar, about five times normal air pressure, so at 1 atm it goes straight from solid to gas at about −78 °C. Its enthalpy of sublimation is 26.1 kJ/mol. The reverse change, gas straight to solid, is deposition. Frost on a car window on a cold, clear night grows this way.