What Is Vapor Pressure? Why Liquids Evaporate
Leave a puddle of water and a puddle of nail-polish remover side by side on a warm day. The acetone is gone in minutes; the water takes hours. Neither one is being heated to its boiling point, and neither one is bubbling. What differs is a property you can put a number on — and the same number explains why you cannot make a decent cup of tea on a mountain.
The short answer: vapour pressure is the pressure exerted by a substance's vapour when it is in dynamic equilibrium with its liquid in a closed container at a given temperature. It measures how strongly a liquid "wants" to escape into the gas phase: weak intermolecular forces and higher temperatures both mean higher vapour pressure.
What vapour pressure actually measures
Half-fill a sealed flask with water and watch what happens at the molecular level.
At first, molecules with above-average energy break free of the surface and enter the space above. As that space fills with vapour, some of those molecules crash back into the liquid and are recaptured. Escape is still happening, and so is return.
Before long the two rates become equal. Molecules keep leaving and keep coming back, but the amount of vapour stops changing. This is dynamic equilibrium — busy, but balanced. The pressure that the vapour exerts at that point is the liquid's vapour pressure at that temperature.
Two features of that definition do a lot of exam work:
- It requires a closed system. In an open room the escaping molecules wander off and never come back, so equilibrium is never reached and the liquid keeps evaporating until it is gone.
- It depends only on the substance and the temperature. Not on how much liquid there is, and not on the surface area. A teaspoon of water and a bathtub of water at 25 °C have exactly the same vapour pressure — surface area changes how fast equilibrium is reached, not where it sits.
Why temperature changes it so dramatically
Raise the temperature and a larger fraction of the molecules have enough kinetic energy to escape the attractions holding them in the liquid. The escape rate climbs, so equilibrium settles at a much higher vapour pressure.
The relationship is not a straight line — it curves upward steeply:
| Temperature | Vapour pressure of water |
|---|---|
| 0 °C | 4.6 mmHg (0.61 kPa) |
| 20 °C | 17.5 mmHg (2.34 kPa) |
| 50 °C | 92.5 mmHg (12.3 kPa) |
| 80 °C | 355 mmHg (47.3 kPa) |
| 100 °C | 760 mmHg (101.3 kPa) |
Note the last row: at 100 °C, water's vapour pressure equals standard atmospheric pressure exactly. That is not a coincidence — it is the definition of water's normal boiling point.
Volatility: which liquids escape easily
A liquid with a high vapour pressure at room temperature is called volatile. Compare a few at 20 °C:
- Diethyl ether — about 440 mmHg. Extremely volatile; boils at 34.6 °C.
- Acetone — about 185 mmHg. Very volatile; boils at 56 °C.
- Ethanol — about 44 mmHg. Moderately volatile; boils at 78 °C.
- Water — about 17.5 mmHg. Not very volatile; boils at 100 °C.
The ordering is the direct consequence of intermolecular forces. Ether is held together by dispersion forces and a weak dipole only — with no O–H it cannot hydrogen-bond to itself at all; water has an extensive hydrogen-bonding network holding every molecule in place. That is why ether escapes so much more easily even though its molecules are four times heavier. Stronger intermolecular forces → lower vapour pressure → higher boiling point. Once you see that chain, half of this topic is done.
Boiling is vapour pressure meeting the atmosphere
A liquid boils when its vapour pressure equals the external pressure pressing down on it. Only then can a bubble of vapour form inside the liquid and survive instead of being crushed.
That single sentence explains a set of results that otherwise look unrelated:
- On Everest, atmospheric pressure is only about 34 kPa (roughly a third of sea level). Water reaches that vapour pressure at about 72 °C, so it boils there — bubbling furiously, but far too cool to brew tea properly or cook rice in a reasonable time.
- In a pressure cooker, the sealed lid lets pressure build to about 2 atm. Water now has to reach roughly 120 °C before its vapour pressure catches up, and food cooks much faster at that higher temperature.
- The "normal boiling point" quoted in tables is the temperature at which vapour pressure equals exactly 1 atm (101.3 kPa). It is a property of the liquid; the actual boiling point depends on where you are standing.
One more useful effect: dissolving a nonvolatile solute (salt, sugar) in a liquid lowers its vapour pressure. It is not that solute particles physically block the surface — the real driver is entropy. Mixing already raises the disorder of the liquid, so escaping into the gas phase gains a solvent molecule less than it would from a pure liquid, and equilibrium settles at a lower pressure. The liquid then needs a higher temperature to reach atmospheric pressure — this is boiling point elevation.
Worked examples
1. Which has a higher vapour pressure at 25 °C, water or ethanol? Ethanol. Its hydrogen-bonding network is weaker (one –OH on a bulky molecule), so more molecules escape and its boiling point is lower.
2. You double the amount of water in a sealed flask at 30 °C. What happens to the vapour pressure? Nothing. Vapour pressure depends on substance and temperature only.
3. Water boils at 95 °C in a kitchen. What does that tell you? The local atmospheric pressure is below 1 atm — the kitchen is at altitude, or a low-pressure weather system has moved in.
4. Why does sweating cool you down? The fastest-moving water molecules escape first, leaving behind a lower-energy — and therefore cooler — liquid. Higher vapour pressure means faster evaporation and faster cooling, which is why rubbing alcohol feels much colder than water on skin.
Common mistakes to avoid
- Thinking a bigger surface area raises vapour pressure. A wide dish evaporates faster, but the equilibrium pressure is identical to that in a narrow tube at the same temperature.
- Believing water can only boil at 100 °C. It boils whenever its vapour pressure matches the surrounding pressure — around 72 °C on Everest, around 120 °C in a pressure cooker.
- Assuming evaporation stops once equilibrium is reached. Molecules never stop leaving. They are just recaptured at the same rate, which is what makes the equilibrium dynamic.
FAQ
What is vapour pressure in simple terms?
It's how hard a liquid's escaped molecules push on the space above it once escape and return have balanced out in a sealed container. High vapour pressure means the liquid evaporates readily.
Does vapour pressure increase with temperature?
Yes, and steeply. More molecules have enough energy to escape at higher temperatures, so the equilibrium vapour pressure rises sharply rather than linearly.
What is the relationship between vapour pressure and boiling point?
They are inversely related. A liquid with high vapour pressure reaches atmospheric pressure at a lower temperature, so it boils sooner. Both are set by the strength of the intermolecular forces.
Why does food cook faster in a pressure cooker?
The sealed pot raises the internal pressure to roughly 2 atm, so water must reach about 120 °C before boiling. The higher temperature speeds up the cooking reactions.
The takeaway
Vapour pressure is the pressure of the vapour sitting in equilibrium above its own liquid, and it is a direct read-out of how weakly the liquid holds itself together. Weak forces and high temperatures push it up; boiling happens the moment it catches up with the pressure of the surrounding air — which is why the boiling point you measure depends on where you measure it.
Pair this with Evaporation vs Boiling: What's the Difference?, What Is a Phase Change? Melting, Boiling, Sublimation, and What Is a Saturated Solution? Solubility Limits. Next in this series: [Boiling Point vs Melting Point].
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