Saturated vapor pressure is the pressure exerted by a vapor when it’s in equilibrium with its liquid (or solid) form in a closed container. At this point, molecules are escaping the liquid surface and returning to it at exactly the same rate, so the amount of vapor holds steady. For water at room temperature (20°C), this pressure is about 2.3 kPa, or roughly 17.5 mmHg. It rises steeply with temperature and is the reason water boils at 100°C at sea level.
How Equilibrium Creates Vapor Pressure
Picture a glass of water sealed inside a container. Molecules at the liquid surface constantly gain enough energy to break free and enter the air above. In an open room, those molecules drift away and the water slowly evaporates. In a sealed container, they can’t escape. They bounce off the walls, collide with each other, and some land back on the liquid surface and get recaptured.
Eventually the number of molecules leaving the surface every second equals the number returning. This balance is called dynamic equilibrium, and the pressure the vapor exerts at that point is the saturated vapor pressure. The vapor is now “saturated” because the air above the liquid physically cannot hold more vapor at that temperature. Add more and it simply condenses back.
Why Temperature Changes Everything
Temperature is the single biggest factor controlling saturated vapor pressure. Warmer molecules move faster and are more likely to break the bonds holding them in the liquid. That means evaporation speeds up. To re-establish equilibrium, more vapor molecules must be present above the surface, which means higher pressure.
The relationship isn’t linear. Saturated vapor pressure increases by roughly 7% for every 1°C (1 K) rise in temperature, so it climbs exponentially. Some concrete numbers for water illustrate the curve:
- 0°C: 0.61 kPa (about 4.6 mmHg)
- 20°C: 2.3 kPa (about 17.5 mmHg)
- 37°C (body temperature): 6.3 kPa (about 47 mmHg)
- 100°C: 101.4 kPa (about 760 mmHg)
Notice that the pressure at 100°C is about 101.4 kPa, which is essentially standard atmospheric pressure. That’s not a coincidence.
The Connection to Boiling
A liquid boils when its saturated vapor pressure equals the surrounding atmospheric pressure. Below that temperature, evaporation happens only at the surface. Once vapor pressure matches the air pressing down on the liquid, bubbles of vapor can form throughout the entire volume of liquid and rise to the surface. That’s boiling.
This is why water boils at a lower temperature on a mountaintop. At high altitude, atmospheric pressure is lower, so the vapor pressure doesn’t need to climb as high to match it. In Denver (about 1,600 meters elevation), water boils near 95°C. In a pressure cooker, the opposite happens: higher pressure inside the sealed pot means water must reach a higher temperature before it can boil, which cooks food faster.
What Makes Some Liquids Evaporate Faster
Different substances have very different saturated vapor pressures at the same temperature, and the reason comes down to how strongly their molecules grip each other. Molecules held together by weak forces escape into the vapor phase easily, producing a high vapor pressure. Molecules bound by strong attractions, like hydrogen bonds, are harder to pry loose, so fewer escape and the vapor pressure stays low.
A clear example: at 25°C, diethyl ether has a vapor pressure of about 520 mmHg, while ethyl alcohol sits at just 75 mmHg. Ether molecules interact through relatively weak forces, so they fly off the surface readily. Alcohol molecules form hydrogen bonds with their neighbors, which takes more energy to break. That’s why ether feels noticeably cold on your skin (it evaporates fast, pulling heat away) while rubbing alcohol evaporates more slowly. Heavier molecules also tend to have lower vapor pressures because their larger electron clouds create stronger intermolecular attractions.
How Dissolved Substances Lower Vapor Pressure
Adding a solute to a liquid, like dissolving salt or sugar in water, lowers the saturated vapor pressure of the solution compared to the pure solvent. The solute molecules occupy space at the liquid surface, physically reducing the number of solvent molecules that can escape. The more solute you add, the greater the drop.
This behavior follows a principle called Raoult’s Law: the vapor pressure of the solvent above a solution is proportional to the fraction of solvent molecules in the mixture. Pure water (100% solvent) has full vapor pressure. A solution that’s 90% water molecules by count has roughly 90% of pure water’s vapor pressure. The identity of the solute barely matters. What counts is how many particles it contributes. Salt, which splits into two ions when dissolved, lowers vapor pressure more per unit than sugar, which stays as a single molecule.
This effect has practical consequences. It’s why saltwater boils at a slightly higher temperature than freshwater, and it’s one reason de-icing salt lowers the freezing point of water on roads.
Saturated Vapor Pressure in Weather
Meteorologists rely on saturated vapor pressure to calculate relative humidity. Relative humidity is simply the actual water vapor pressure in the air divided by the saturated vapor pressure at that air temperature, expressed as a percentage. When the air holds 50% of the vapor it could hold at that temperature, relative humidity is 50%.
Because saturated vapor pressure rises so steeply with temperature, warm air can hold far more moisture than cold air. Air at 30°C can hold roughly four times the water vapor that air at 10°C can. This is why humid summer days feel so oppressive: the actual amount of water vapor in the air is high. It’s also why cooling air leads to condensation. As temperature drops, the saturated vapor pressure falls until it meets the actual vapor pressure. At that point the air is at 100% humidity, and dew, fog, or clouds begin to form. The temperature where this happens is called the dew point.
Calculating Saturated Vapor Pressure
For quick estimates, scientists and engineers commonly use the Antoine equation:
log₁₀(P) = A − B / (T + C)
Here, P is the vapor pressure (typically in bar), T is the temperature (in Kelvin), and A, B, and C are constants specific to each substance. For ethanol between about 0°C and 79°C, the constants are A = 5.37, B = 1670.4, and C = −40.2, as published by NIST. Different temperature ranges use slightly different constants because the equation is an approximation. Water, ethanol, and hundreds of other compounds have published Antoine constants you can look up in chemistry reference databases.
A more precise approach uses the Clausius-Clapeyron equation, which derives vapor pressure from the energy required to vaporize a substance. For most everyday and engineering purposes, though, the Antoine equation or published steam tables give accurate enough results without heavy math.
Phase Diagrams and the Critical Point
On a phase diagram, which plots pressure against temperature, the saturated vapor pressure curve is the boundary line between liquid and gas. Any combination of pressure and temperature that falls on this curve means liquid and vapor coexist in equilibrium. Move to one side and you have only liquid; move to the other and you have only gas.
This boundary doesn’t extend forever. It ends at a point called the critical point, where the distinction between liquid and gas disappears. For water, the critical point is at about 374°C and 22.1 MPa. Above those conditions, water exists as a supercritical fluid that has properties of both liquid and gas simultaneously. Below the curve, there’s also a boundary between solid and vapor (sublimation), which is why ice can slowly disappear in a freezer even without melting first.