Vapour, Evaporation & Boiling
When attraction wins: liquefaction, evaporation, and why water boils at 100°C only at sea level.
Lesson at a glanceShow the key points
- Compression or cooling liquefies a gas, except above the critical temperature (water 374.1°C). A gas below it is a vapour.
- Isotherms flatten where vapour and liquid coexist (e.g. water at 350°C liquefies at 163 atm).
- Evaporation: fast surface molecules escape; it cools the remaining liquid; wind speeds it up.
- Saturated vapour: maximum amount a space can hold; its pressure (SVP) depends only on temperature and rises steeply with it.
- Boiling: vapour forms throughout the liquid when SVP equals the external pressure (100°C at 1 atm, 82°C at 0.5 atm).
Kinetic theory ignores molecular attraction. Squeeze or cool a gas enough and attraction takes over: the gas condenses. This is the physics of vapour, evaporation and boiling.
Vapour and the critical temperature
At low pressure and high temperature molecules are far apart and their attraction is weak compared with their kinetic energy. Compress the gas or cool it and the attraction grows in importance, until the gas liquefies. But above a certain temperature the molecules are so energetic that no pressure can liquefy the gas.
On a pressure–volume graph each curve at fixed temperature is an isotherm. Water vapour at 350°C begins to liquefy at about 163 atm; at 370°C, about 194 atm. Along the flat part vapour and liquid coexist; once all is liquid, the curve rises steeply because liquids barely compress. Above 374.1°C the isotherm never flattens.
Evaporation
Leave ether in three dishes, one with a lid, one under a large jar, one open. After a few hours the lidded dish is nearly full, the jar has lost some, the open dish is almost empty. Molecules near the surface occasionally escape (evaporation). In a closed space they pile up until as many return to the liquid each second as leave it; in the open they drift away, and wind carries them off even faster, which is why clothes dry quicker on a windy day.
Saturated vapour
In a closed space above a liquid, equilibrium is reached when evaporation and return balance. Inject more vapour and the extra condenses; take some away and more evaporates. The space can hold only a certain maximum amount of vapour at a given temperature.
The space holds the maximum possible vapour. Its pressure is the saturation vapour pressure (SVP), fixed by the temperature.
Less than the maximum; more liquid will evaporate if available.
Higher temperature means faster molecules and more escapes, so the SVP rises steeply with temperature:
| °C | SVP of water (mm Hg) | °C | SVP (mm Hg) |
|---|---|---|---|
| 0 | 4.58 | 80 | 355 |
| 10 | 8.94 | 100 | 760 |
| 20 | 17.5 | 120 | 1490 |
| 40 | 55.1 | 140 | 2710 |
| 60 | 149 | 160 | 4630 |
Saturation doesn’t need a liquid present: keep injecting vapour into a closed vessel and condensation starts once the SVP is reached. Air holding the maximum vapour is “saturated”.
Boiling
Heat a liquid and its molecules gain energy everywhere, not just at the surface. Eventually bubbles of vapour form throughout and rise out: boiling. Bubbles can survive only when the vapour pressure inside them can match the pressure on the liquid:
What external pressure makes water boil at 120°C? Use the table.
Evaporation versus boiling
Evaporation: only fast molecules at the surface escape, at any temperature. Boiling: molecules throughout the liquid form vapour, at one temperature for a given pressure.
- 1Above the critical temperature (water: 374.1°C) a gas can't be liquefied by pressure; below it, it is a vapour.
- 2Evaporation: fast surface molecules escape; it cools the liquid; wind speeds it up.
- 3Saturated vapour: the maximum the space can hold; its pressure (SVP) depends only on temperature and rises steeply.
- 4Boiling: when SVP equals the external pressure.
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