Vapor Pressure

Vapor pressure is the partial pressure exerted by gas molecules in equilibrium with a liquid in a closed system. At dynamic equilibrium, vaporization and condensation occur at equal rates, with molecules continuously exchanging between phases but no net change in liquid or vapor mass. This measurable quantity reflects the tendency of liquid molecules to escape into the vapor phase at a given temperature.

Strong intermolecular forces impede vaporization and favor recapture of gas molecules, resulting in low vapor pressure. Weak intermolecular attractions present less barrier to vaporization, yielding high vapor pressures. For example, diethyl ether has weak dispersion forces and high vapor pressure, while water exhibits strong hydrogen bonding and lower vapor pressure.

Heating raises the liquid's temperature, increasing molecular thermal energies. At higher temperatures, a greater fraction of molecules possess sufficient energy to overcome intermolecular forces and escape into the vapor phase. Both the increased escape rate and greater average speed of escaping molecules contribute to higher vapor pressure.

Boiling occurs when a liquid's vapor pressure equals the external atmospheric pressure. The normal boiling point is the temperature at which vapor pressure reaches 1 atm. At higher external pressures, more vapor molecules are needed to balance the pressure, so boiling occurs at higher temperatures. Conversely, at lower pressures, boiling happens at lower temperatures.

Volatility qualitatively describes the tendency of liquid molecules to escape into the vapor phase, while vapor pressure is the quantitative measure of this tendency under specific conditions. Volatility is a comparative term based on vapor pressures of liquids held under identical conditions. For instance, hexane is volatile relative to water because hexane has weaker intermolecular forces and higher vapor pressure.

At higher altitudes, atmospheric pressure is lower than 1 atm. Water's vapor pressure requires fewer molecules to equal this reduced external pressure, so boiling occurs at lower temperatures. Conversely, in a pressure cooker with higher external pressure, water must reach higher temperatures before its vapor pressure equals the elevated external pressure.

Neither the surface area of the liquid in contact with vapor nor the size of the vessel affects the vapor pressure at a given temperature. However, these factors do influence the time required for vapor-liquid equilibrium to be established. Vapor pressure depends only on the chemical identity of the substance and temperature.