Vapor pressure

Vapor pressure is measured in the standard units of pressure. The International System of Units (SI) recognizes pressure as a derived unit with the dimension of force per area and designates the pascal (Pa) as its standard unit. The most accurate results are obtained near the boiling point of the substance; measurements smaller than are subject to major errors. Procedures often consist of purifying the test substance, isolating it in a container, evacuating any foreign gas, then measuring the equilibrium pressure of the gaseous phase of the substance in the container at different temperatures. Better accuracy is achieved when care is taken to ensure that the entire substance and its vapor are both at the prescribed temperature. This is often done, as with the use of an isoteniscope, by submerging the containment area in a liquid bath. Very low vapor pressures of solids can be measured using the Knudsen effusion cell method. In a medical context, vapor pressure is sometimes expressed in other units, specifically millimeters of mercury (mmHg). Accurate knowledge of the vapor pressure is important for volatile inhalational anesthetics, most of which are liquids at body temperature but have a relatively high vapor pressure. ==Estimating vapor pressures with Antoine equation==

The Antoine equation is a pragmatic mathematical expression of the relation between the vapor pressure and the temperature of pure liquid or solid substances. It is obtained by curve-fitting and is adapted to the fact that vapor pressure is usually increasing and concave as a function of temperature. The basic form of the equation is: :\log P = A-\frac{B}{C+T} and it can be transformed into this temperature-explicit form: :T = \frac{B}{A-\log P} - C where: • P is the absolute vapor pressure of a substance • T is the temperature of the substance • A, B and C are substance-specific coefficients (i.e., constants or parameters) • \log is typically either \log_{10} or \log_e gives "one of the best" fits to experimental data but is quite complex. It expresses reduced vapor pressure as a function of reduced temperature. ==Relation to boiling point of liquids==

As a general trend, vapor pressures of liquids at ambient temperatures increase with decreasing boiling points. This is illustrated in the vapor pressure chart (see right) that shows graphs of the vapor pressures versus temperatures for a variety of liquids. At the normal boiling point of a liquid, the vapor pressure is equal to the standard atmospheric pressure defined as 1 atmosphere, 760Torr, 101.325kPa, or 14.69595psi. For example, at any given temperature, methyl chloride has the highest vapor pressure of any of the liquids in the chart. It also has the lowest normal boiling point at , which is where the vapor pressure curve of methyl chloride (the blue line) intersects the horizontal pressure line of one atmosphere (atm) of absolute vapor pressure. Although the relation between vapor pressure and temperature is non-linear, the chart uses a logarithmic vertical axis to produce slightly curved lines, so one chart can graph many liquids. A nearly straight line is obtained when the logarithm of the vapor pressure is plotted against 1/(T + 230) where T is the temperature in degrees Celsius. The vapor pressure of a liquid at its boiling point equals the pressure of its surrounding environment. ==Liquid mixtures: Raoult's law==

Raoult's law gives an approximation to the vapor pressure of mixtures of liquids. It states that the activity (pressure or fugacity) of a single-phase mixture is equal to the mole-fraction-weighted sum of the components' vapor pressures: : P_{\rm tot} =\sum_i P y_i = \sum_i P_i^{\rm sat} x_i \, where P_{\rm tot} is the mixture's vapor pressure, x_i is the mole fraction of component i in the liquid phase and y_i is the mole fraction of component i in the vapor phase respectively. P_i^{\rm sat} is the vapor pressure of component i. Raoult's law is applicable only to non-electrolytes (uncharged species); it is most appropriate for non-polar molecules with only weak intermolecular attractions (such as London forces). Systems that have vapor pressures higher than indicated by the above formula are said to have positive deviations. Such a deviation suggests weaker intermolecular attraction than in the pure components, so that the molecules can be thought of as being "held in" the liquid phase less strongly than in the pure liquid. An example is the azeotrope of approximately 95% ethanol and water. Because the azeotrope's vapor pressure is higher than predicted by Raoult's law, it boils at a temperature below that of either pure component. There are also systems with negative deviations that have vapor pressures that are lower than expected. Such a deviation is evidence for stronger intermolecular attraction between the constituents of the mixture than exists in the pure components. Thus, the molecules are "held in" the liquid more strongly when a second molecule is present. An example is a mixture of trichloromethane (chloroform) and 2-propanone (acetone), which boils above the boiling point of either pure component. The negative and positive deviations can be used to determine thermodynamic activity coefficients of the components of mixtures. ==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 dissolution to form a vapor. For most solids this pressure is very low, but some notable exceptions are naphthalene, dry ice (the vapor pressure of dry ice is 5.73 MPa (831 psi, 56.5 atm) at 20 °C, which causes most sealed containers to rupture), and ice. 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. There are a number of methods for calculating the sublimation pressure (i.e., the vapor pressure) of a solid. One method is to estimate the sublimation pressure from extrapolated liquid vapor pressures (of the supercooled liquid), if the heat of fusion is known, by using this particular form of the Clausius–Clapeyron relation: :\ln\,P^{\rm sub}_{\rm s} = \ln\,P^{\rm sub}_{\rm l} - \frac{\Delta_{\rm fus}H}{R} \left( \frac{1}{T_{\rm sub}} - \frac{1}{T_{\rm fus}} \right) where: • P^{\rm sub}_{\rm s} is the sublimation pressure of the solid component at the temperature T_{\rm sub} . • P^{\rm sub}_{\rm l} is the extrapolated vapor pressure of the liquid component at the temperature T_{\rm sub} . • \Delta_{\rm fus}H is the heat of fusion. • R is the gas constant. • T_{\rm sub} is the sublimation temperature. • T_{\rm fus} is the melting point temperature. This method assumes that the heat of fusion is temperature-independent, ignores additional transition temperatures between different solid phases, and it gives a fair estimation for temperatures not too far from the melting point. It also shows that the sublimation pressure is lower than the extrapolated liquid vapor pressure (ΔfusH > 0) and the difference grows with increased distance from the melting point. ==Boiling point of water==

or 101.325kPa. Like all liquids, water boils when its vapor pressure reaches its surrounding pressure. In nature, the atmospheric pressure is lower at higher elevations and water boils at a lower temperature. The boiling temperature of water for atmospheric pressures can be approximated by the Antoine equation: :\log_{10}\left(\frac{P}{1\text{ Torr}}\right) = 8.07131 - \frac{1730.63\ {}^\circ\text{C}}{233.426\ {}^\circ\text{C} + T_b} or transformed into this temperature-explicit form: :T_b = \frac{1730.63\ {}^\circ\text{C}}{8.07131 - \log_{10}\left(\frac{P}{1\text{ Torr}}\right)} - 233.426\ {}^\circ\text{C} where the temperature T_b is the boiling point in degrees Celsius and the pressure P is in torr. ==Dühring's rule==

Dühring's rule states that a linear relationship exists between the temperatures at which two solutions exert the same vapor pressure. ==Examples==

The following table is a list of a variety of substances ordered by increasing vapor pressure (in absolute units). ==Estimating vapor pressure from molecular structure==

Several empirical methods exist to estimate the vapor pressure from molecular structure for organic molecules. Some examples are SIMPOL.1 method, the method of Moller et al., ==Meaning in meteorology==

In meteorology, the term vapor pressure means the partial pressure of water vapor in the atmosphere, even if it is not in equilibrium. This differs from its meaning in other sciences. Relative humidity is defined relative to saturation vapor pressure. The still-current term saturation vapor pressure derives from the obsolete theory that water vapor dissolves into air, and that air at a given temperature can only hold a certain amount of water before becoming "saturated". ==See also==