The vapor pressure of a fluid is a crucial concept in fluid mechanics, influencing phenomena such as boiling and cavitation. Vapor pressure refers to the pressure exerted by a vapor at a state of thermodynamic equilibrium with its corresponding liquid phase at a specific temperature. It represents the tendency of molecules to escape from the fluid surface into the vapor phase.

When a liquid is placed in a closed container with a small air space, and the space is evacuated, vapor molecules will escape from the liquid into the space. As more molecules escape, the pressure in the space increases until an equilibrium is reached, where the rate of molecules escaping the liquid equals the rate of molecules returning to it. This equilibrium pressure is the vapor pressure of the liquid.

Vapor pressure is temperature-dependent. As the temperature increases, the molecules' kinetic energy increases, resulting in more molecules overcoming intermolecular forces and escaping into the vapor phase. It follows that the vapor pressure increases with temperature. For instance, water at 100°C has a vapor pressure of 101.3 kPa (1 atm), allowing it to boil at standard atmospheric pressure. However, at higher elevations, where atmospheric pressure is lower, water boils at lower temperatures because its vapor pressure equals the reduced atmospheric pressure at the higher elevation.