Vapor Pressure
Definition and meaning of Vapor Pressure in chemistry.
Vapor pressure is the pressure exerted by a vapor in equilibrium with its liquid (or solid) phase at a given temperature. It is unique to each substance and increases predictably with temperature.
In more detail
At any temperature, molecules in a liquid constantly evaporate while gas molecules simultaneously condense back, eventually reaching equilibrium where both processes occur at equal rates. The pressure exerted by the gas phase at this equilibrium is the vapor pressure. Substances with weaker intermolecular forces (such as diethyl ether) have higher vapor pressures than those with strong forces (such as water with hydrogen bonding). A liquid boils when its vapor pressure equals the atmospheric pressure above it.
Key facts
| Definition | Pressure exerted by a vapor in equilibrium with its condensed phase |
|---|---|
| Temperature dependence | Increases exponentially with increasing temperature |
| Field | Physical Chemistry |
| Boiling point connection | Liquid boils when vapor pressure equals atmospheric pressure |
Water has a vapor pressure of approximately 23.8 mmHg at 25°C and 760 mmHg (1 atmosphere) at 100°C. This is why water boils at 100°C at sea level, where atmospheric pressure is 1 atm.
Frequently asked questions
What is the relationship between vapor pressure and boiling point?
A liquid boils when its vapor pressure equals the atmospheric pressure above it. At higher altitudes where atmospheric pressure is lower, liquids boil at lower temperatures because they reach lower vapor pressures.
Why do some liquids evaporate faster than others?
Liquids with weaker intermolecular forces have higher vapor pressures, meaning more molecules escape to the gas phase. These volatile liquids evaporate faster because a larger fraction of molecules have sufficient energy to overcome intermolecular attractions.