Vapor Pressure
Definition and meaning of Vapor Pressure in chemistry.
Vapor pressure is the pressure exerted by a gas sitting directly above its liquid. This gas must be in a state of dynamic equilibrium with the liquid phase. Every unique chemical substance has a specific vapor pressure that increases as it warms.
In more detail
At any given temperature, molecules in a liquid are constantly moving and colliding. Some molecules near the surface gain enough energy to escape into the gas phase. We call this continuous escaping process by the common name of evaporation.
At the exact same time, some gas molecules lose energy and return to the liquid. We call this returning process by the common name of condensation. When evaporation and condensation happen at the exact same rate, the system reaches equilibrium.
The pressure created by the gas at this balanced stage is the vapor pressure. Substances with very weak intermolecular forces let molecules escape much more easily. These volatile substances have high vapor pressures and evaporate very quickly in open air.
Substances with strong intermolecular forces hold their molecules tightly and have low vapor pressures. As you heat any liquid, more molecules gain the kinetic energy needed to escape. This means that vapor pressure always goes up predictably as the temperature goes up.
A liquid finally boils when its vapor pressure equals the surrounding atmospheric pressure. Students often confuse a liquid's vapor pressure with the pressure of the outside air. Vapor pressure only comes from the evaporated gas molecules of the substance itself.
Key facts
| Definition | Pressure of a gas in equilibrium with its liquid |
|---|---|
| Temperature Dependence | Always increases as the temperature increases |
| Field | Physical Chemistry |
| Intermolecular Forces | Weaker forces result in a higher vapor pressure |
| Boiling Point Condition | Liquid boils when vapor pressure equals atmospheric pressure |
| Equilibrium State | Evaporation rate exactly equals the condensation rate |
Water has a relatively low vapor pressure because of its strong hydrogen bonds. At room temperature, the vapor pressure of water is only about 24 millimeters of mercury. If you heat the water to 100 degrees Celsius, the vapor pressure increases dramatically. It reaches exactly 760 millimeters of mercury, which equals standard atmospheric pressure at sea level. Because its vapor pressure now matches the outside air pressure, the water begins to boil.
Frequently asked questions
What is the relationship between vapor pressure and boiling point?
A liquid boils when its internal vapor pressure matches the outside atmospheric pressure. If you go to a high mountain, water boils at a much lower temperature.
Why do some liquids evaporate much faster than others?
Liquids with weak intermolecular forces do not hold their molecules together very tightly. Many molecules escape into the gas phase, creating a very high vapor pressure.
Does the size of the container affect the vapor pressure?
No, the size of the container does not change the equilibrium vapor pressure. It only depends on the specific chemical substance and the current temperature.