n.
The pressure exerted by a vapor in equilibrium with its solid or liquid phase.
vapor pressure
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n.
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The saturation pressures exerted by vapors which are in equilibrium with their liquid or solid forms. One of the most important physical properties of a liquid, the vapor pressure, enters into many thermodynamic calculations and underlies several methods for the determination of the molecular weights of substances dissolved in liquids. For a discussion of the vapor pressure relationships of solids see Sublimation. See also Molecular weight; Solution.
If a liquid is introduced into an evacuated vessel at a given temperature, some of the liquid will vaporize, and the pressure of the vapor will attain a maximum value which is termed the vapor pressure of the liquid at that temperature. Although the quantity of liquid remaining does not diminish thereafter, the process of evaporation does not cease. A dynamic equilibrium is established, in which molecules escape from the liquid phase and return from the vapor phase at equal rates. See also Evaporation.
It is important to make a distinction between the vapor pressure of a liquid, as described above, and the pressure of a vapor. The vapor pressure of a pure liquid is a unique and characteristic property of the liquid and depends only upon the temperature. A gas or vapor may, on the other hand, exert any pressure within reason, depending upon the volume to which it is confined, provided it is not in contact with its liquid phase. See also Phase equilibrium.
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The component of the total pressure which is caused by the presence of a vapor, as, for example, by the presence of water vapor in air.
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vapor pressure, pressure exerted by a vapor that is in equilibrium with its liquid. A liquid standing in a sealed beaker is actually a dynamic system: some molecules of the liquid are evaporating to form vapor and some molecules of vapor are condensing to form liquid. At equilibrium the rates of the two processes are equal and the system appears to be stationary (see chemical equilibrium). The vapor, like any gas, exerts a pressure, and this pressure at equilibrium is called the vapor pressure. Vapor pressure depends on various factors, the most important of which is the nature of the liquid. If the molecules of liquid bind to each other very strongly, there will be less tendency for the molecules to escape as gas and a consequent lower vapor pressure; for example, polar molecules that can form hydrogen bonds between themselves, e.g., water molecules and the alcohols, have relatively low vapor pressures. If there is only weak interaction between the liquid molecules, there will be a greater tendency for the molecules to evaporate and a higher vapor pressure. Temperature also affects the vapor pressure. If the system in equilibrium is perturbed by raising the temperature, then according to Le Châtelier's principle the system should react to relieve this stress; as the temperature is increased, the evaporation process, which absorbs heat, is speeded up to a greater degree than the condensation process, which gives off heat, so that the vapor pressure is higher when equilibrium is restored at the new temperature. If the temperature is increased enough to raise the vapor pressure until it equals atmospheric pressure, the liquid will boil. If the external pressure is reduced, as in a vacuum system, then the liquid will boil much more readily than under atmospheric pressure. This fact is used in the vacuum distillation process to obtain relatively pure samples of liquids with high boiling points. Some solids, e.g., iodine and carbon dioxide, are capable of subliming (going directly from a solid to a gas) at atmospheric pressure and room temperature; thus, such solids also have significant vapor pressures under these conditions. Another factor affecting vapor pressure is the presence of dissolved substances in the liquid or solid; according to Raoult's law, the vapor pressure of a pure liquid or solid is lowered by the addition of a solute.
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In physics and
| Wikipedia: Vapor pressure |
Vapor pressure (also known as equilibrium vapor pressure), is the pressure of a vapor in equilibrium with its non-vapor phases. All liquids and solids have a tendency to evaporate to a gaseous form, and all gases have a tendency to condense back into their original form (either liquid or solid). At any given temperature, for a particular substance, there is a pressure at which the gas of that substance is in dynamic equilibrium with its liquid or solid forms. This is the vapor pressure of that substance at that temperature. The equilibrium vapor pressure is an indication of a liquid's evaporation rate. It relates to the tendency of molecules and atoms to escape from a liquid or a solid. A substance with a high vapor pressure at normal temperatures is often referred to as volatile.
The vapor pressure of any substance increases non-linearly with temperature according to the Clausius-Clapeyron relation. The atmospheric pressure boiling point of a liquid (also known as the normal boiling point) is the temperature where the vapor pressure equals the ambient atmospheric pressure. With any incremental increase in that temperature, the vapor pressure becomes sufficient to overcome atmospheric pressure and lift the liquid to form bubbles inside the bulk of the substance. Bubble formation deeper in the liquid requires a higher pressure, and therefore higher temperature, because the fluid pressure increases above the atmospheric pressure as the depth increases.
Contents |
Relation between vapor pressures and normal boiling points of liquids
The higher the vapor pressure of a liquid at a given temperature, the lower the normal boiling point (i.e., the boiling point at atmospheric pressure) of the liquid.
The vapor pressure chart to the right has graphs of the vapor pressures versus temperatures for a variety of liquids.[1] As can be seen in the chart, the liquids with the highest vapor pressures have the lowest normal boiling points.
For example, at any given temperature, propane has the highest vapor pressure of any of the liquids in the chart. It also has the lowest normal boiling point (−42.1 °C), which is where the vapor pressure curve of propane (the purple 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 in order to obtain slightly curved lines so that one chart can graph many liquids.
Units
The international SI unit for pressure is the pascal (Pa), equal to one newton per square meter (N·m-2 or kg·m-1·s-2). The conversions to other pressure units are:
pascal (Pa) |
bar (bar) |
technical atmosphere (at) |
atmosphere (atm) |
torr (Torr) |
pound-force per square inch (psi) |
|
|---|---|---|---|---|---|---|
| 1 Pa | ≡ 1 N/m2 | 10−5 | 1.0197×10−5 | 9.8692×10−6 | 7.5006×10−3 | 145.04×10−6 |
| 1 bar | 100,000 | ≡ 106 dyn/cm2 | 1.0197 | 0.98692 | 750.06 | 14.5037744 |
| 1 at | 98,066.5 | 0.980665 | ≡ 1 kgf/cm2 | 0.96784 | 735.56 | 14.223 |
| 1 atm | 101,325 | 1.01325 | 1.0332 | ≡ 1 atm | 760 | 14.696 |
| 1 torr | 133.322 | 1.3332×10−3 | 1.3595×10−3 | 1.3158×10−3 | ≡ 1 Torr; ≈ 1 mmHg | 19.337×10−3 |
| 1 psi | 6,894.76 | 68.948×10−3 | 70.307×10−3 | 68.046×10−3 | 51.715 | ≡ 1 lbf/in2 |
Example reading: 1 Pa = 1 N/m2 = 10−5 bar = 10.197×10−6 at = 9.8692×10−6 atm, etc.
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 5.73 MPa (831 psi, 56.5 atm) at 20 degrees Celsius, meaning it will cause most sealed containers to explode), 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.
The sublimation pressure can be calculated[2] from extrapolated liquid vapor pressures (of the supercooled liquid) if the heat of fusion is known. The heat of fusion has to be added in addition to the heat of vaporization to evaporize a solid. Assuming that the heat of fusion is temperature-independent and ignoring additional transition temperatures between different solid phases the equation
with:
 |
= Sublimation pressure of the solid component at the temperature T < Tm |
|---|---|
 |
= Extrapolated vapor pressure of the liquid component at the temperature T < Tm |
| ΔHm | = Heat of fusion |
| R | = Gas constant |
| T | = Sublimation temperature |
| Tm | = Melting point temperature |
gives a fair estimation for temperatures not too far from the melting point. This equation also shows that the sublimation pressure is lower than the extrapolated liquid vapor pressure (ΔH m is positive) and the difference grows with increased distance from 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 atmospheric pressure can be approximated by this Antoine equation:
or transformed into this temperature-explicit form:
where the temperature Tb is the boiling point temperature in degrees Celsius and the pressure P is in Torr.
Mixtures
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:
| ptot = | ∑ | piχi |
| i |
where p is vapor pressure, i is a component index, and χ is a mole fraction. The term piχi is the vapor pressure of component i in the mixture. 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.
Usage of the term "vapor pressure" in meteorology
In meteorology, the term vapor pressure is used to mean the partial pressure of water vapor in the atmosphere, even if it is not equilibrium,[3] and the equilibrium vapor pressure is specified as such. Meteorologists also use the term saturation vapor pressure to refer to the equilibrium vapor pressure of water or brine above a flat surface, to distinguish it from equilibrium vapor pressure which takes into account the shape and size of water droplets and particulates in the atmosphere.[4]
See also
- Absolute humidity
- Clausius-Clapeyron Equation
- Partial pressure
- Relative humidity
- Relative volatility
- Triple point
- Vapor-liquid equilibrium
- Vapor Pressure of Water at Various Temperatures
- Volatility
- Antoine equation
References
- ^ Perry, R.H. and Green, D.W. (Editors) (1997). Perry's Chemical Engineers' Handbook (7th Edition ed.). McGraw-Hill. ISBN 0-07-049841-5.
- ^ Moller B., Rarey J., Ramjugernath D., "Estimation of the vapour pressure of non-electrolyte organic compounds via group contributions and group interactions ", J.Mol.Liq., 143(1), 52-63, 2008
- ^ Glossary (Developed by the American Meteorological Society)
- ^ A Brief Tutorial (An article about the definition of equilibrium vapor pressure)
External links
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 | Dictionary. The American Heritage® Dictionary of the English Language, Fourth Edition Copyright © 2007, 2000 by Houghton Mifflin Company. Updated in 2007. Published by Houghton Mifflin Company. All rights reserved. Read more |
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Mentioned in
- sublimation pressure (thermodynamics)
- condensed system (physical chemistry)
- sublimation curve (thermodynamics)
- Raoult's law (physical chemistry)
- relative humidity (meteorology)
- boiling point
- equilibrium vapor pressure (physics)
- psychrometric formula (thermodynamics)
- Antoine equation (physics)
- Ramsay-Young rule (thermodynamics)
- saturation deficit (meteorology)
- Cox chart (chemistry)
- Dühring's rule (physical chemistry)
- saturated air (meteorology)
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