Vapor pressure

Vapor pressure is the pressure of a vapor in equilibrium with its non-vapor phases. Most often the term is used to describe a liquid's tendency to evaporate. It is the tendency of molecules and atoms to escape from a liquid or a solid. At any given temperature, for a particular substance, there is a pressure at which the vapor of that substance is in equilibrium with its liquid or solid forms. This is the equilibrium vapor pressure or saturation vapor pressure of that substance at that temperature. The term vapor pressure is often understood to mean the saturation vapor pressure. A substance with a high vapor pressure at normal temperatures is often referred to as volatile. The higher the vapor pressure of a material at a given temperature, the lower the boiling point.

Units of vapor pressure

The most common unit for vapor pressure is the torr. 1 torr = 1 mm Hg (one millimeter of mercury). The international unit for pressure is: 1 pascal = a force of 1 newton per square meter = 10 dyn/cm² = 0.01 mbar= 0.0075 mmHg = 0.00000969 atm= 0.00014 psi .

Equilibrium vapor pressure of solids

Equilibrium vapor pressure can be defined as the pressure reached when a condensed phase is in equilibrium with its own vapor. In the case of an equilibrium solid, such as a crystal, this can be defined as the pressure when the rate of sublimation of a solid matches the rate of deposition of its vapor phase. For most solids this pressure is very low, but some notable exceptions are naphthalene, dry ice (the vapor pressure of dry ice is 831psi (56.55atm) at 20 degrees Celsius, meaning it will cause most non-ventilated containers to explode if sealed inside).and ice. Ice will still continue to disappear even though the ambient temperature is below the freezing point of water. All solid materials have a vapor pressure. However, due to their often extremely low values, measurement can be rather difficult. Typical techniques include the use of thermogravimetry and gas transpiration.

Relation between solid and liquid vapor pressures

It may be noted that the vapor pressure of a substance in liquid form is usually different from the vapor pressure of the same substance in solid form. If the temperature is such that the vapor pressure of the liquid is higher than that of the solid, liquid will vaporize but vapor will condense to a solid, i.e. the liquid is freezing. If the temperature is such that the vapor pressure of the liquid is lower than that of the solid, solid will vaporize but vapor will condense to a liquid, i.e. the solid is melting. At the temperature that equalizes the two vapor pressures, an equilibrium exists between solid and liquid phases. This temperature is referred to as the melting point.

Water vapor pressure

Water, like all liquids, starts to boil when its vapor pressure reaches its surrounding pressure. At higher elevations the atmospheric pressure is lower and water will boil at a lower temperature. The boiling temperature of water for pressures around 100 kPa can be approximated by

$T_{v}=100+0.0002772\cdot (p-101000)-1.24\cdot 10^{-9}\cdot (p-101000)^{2}$

where the temperature is in degrees Celsius and the pressure p is in pascals. One gets the vapor pressure by solving this equation for p.

Partial pressures

Raoult's law approximately governs the vapor pressure of mixtures of liquids.

Examples of vapor pressures

gas	vapour pressure (bar)	vapour pressure (mmHg)	notes
Helium	1	750	@ -269.15 °C
Propane	22	16500	@ 55 °C
Butane	2.2	1650	@ 20 °C
Carbonyl sulfide	12.55	9412	@ 25 °C
Acetaldehyde	0.987	740	@ 20 °C
Freon 113	0.379	284	@ 20 °C
Methyl isobutyl ketone	0.02648	19.86	@ 25 °C
Tungsten	0.001	0.75	@ 3203 °C

External links