Vaporpressure

Vapor pressure is the pressure exerted by a vapor in thermodynamic equilibrium with its condensed phase (liquid or solid) at a given temperature. It represents the escaping tendency of molecules from the condensed phase into the vapor. For a pure substance, the saturated vapor pressure is the equilibrium vapor pressure at that temperature.

Vapor pressure varies with temperature: as temperature increases, more molecules have sufficient energy to overcome intermolecular forces, so the vapor pressure rises. Substances with relatively high vapor pressures at ambient temperatures are considered volatile.

Thermodynamic relations describe the temperature dependence. The Clausius-Clapeyron equation relates changes in vapor pressure to the enthalpy of vaporization: d(ln P)/dT ≈ ΔHvap/(R T^2). Integrating gives forms such as ln P = -ΔHvap/(R T) + constant; the Antoine equation provides a practical empirical correlation P = f(T) using substance-specific constants.

Vapor pressure is typically expressed in units of pascals (Pa) or kilopascals (kPa); other common units include Torr or millimeters of mercury (mmHg). It can be measured by static methods that equilibrate a liquid in a sealed vessel or by dynamic methods that monitor vapor flow or condensation.

In mixtures, each component has a partial vapor pressure. Under Raoult's law for ideal solutions, P_i = x_i P_i*, where x_i is the mole fraction and P_i* is the pure-component vapor pressure. The total vapor pressure is the sum of the partial pressures.

Vapor pressure influences boiling points, distillation, evaporation rates, and humidity calculations. It is central to meteorology, chemical engineering, and materials science, affecting storage, drying processes, and design of thermal systems.