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| [[File:Vapor pressure.svg|thumb|The picture shows the particle transition, as a result of their vapor pressure, from the liquid phase to the gas phase and converse.]]
| | My name is Maryanne (44 years old) and my hobbies are Sand castle building and Aircraft spotting.<br><br>Stop by my homepage - [http://www.Rogercarthew.com/blog/?p=2196& Fifa 15 Coin Generator] |
| '''Vapor pressure''' or '''equilibrium vapor pressure''' is defined as the [[pressure]] exerted by a [[vapor]] in [[thermodynamic equilibrium]] with its [[Condensation|condensed]] [[Phase (matter)|phase]]s (solid or liquid) at a given temperature in a [[Thermodynamic_system#Closed_system|closed system]]. The equilibrium vapor pressure is an indication of a liquid's [[evaporation]] rate. It relates to the tendency of particles to escape from the liquid (or a solid). A substance with a high vapor pressure at normal temperatures is often referred to as ''[[volatility (chemistry)|volatile]]''.
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| 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 at which 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 vapor bubbles inside the bulk of the substance. [[liquid bubble|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.
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| The vapor pressure that a single component in a mixture contributes to the total pressure in the system is called [[partial pressure]]. For example, air at sea level, and saturated with water vapor at 20 °C, has partial pressures of about 23 mbar of [[water]], 780 mbar of [[nitrogen]], 210 mbar of [[oxygen]] and 9 mbar of [[argon]].
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| ==Measurement and units==
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| Vapor pressure is measured in the standard units of [[pressure]]. The [[International System of Units]] (SI) recognizes pressure as a [[SI derived unit|derived unit]] with the dimension of force per area and designates the [[pascal (unit)|pascal]] (Pa) as its standard unit. One pascal is one [[newton (unit)|newton]] per [[square meter]] (N·m<sup>−2</sup> or kg·m<sup>−1</sup>·s<sup>−2</sup>).
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| Experimental measurement of vapor pressure is a simple procedure for common pressures between 1 and 200 kPa.<ref>{{cite web|url=http://www.capec.kt.dtu.dk/documents/overview/Vapor-pressure-Ruzicka.pdf|title=Vapor Pressure of Organic Compounds. Measurement and Correlation|author=K. Růžička, M. Fulem, V. Růžička}}</ref> Most accurate results are obtained near the boiling point of substances and large errors result for measurements smaller than {{gaps|1|kPa}}. 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 at the prescribed temperature. This is often done, as with the use of an [[isoteniscope]], by submerging the containment area in a liquid bath.
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| ==Estimating vapor pressures with Antoine equation==
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| The [[Antoine equation]] <ref name=frostburg>[http://antoine.frostburg.edu/chem/senese/101/liquids/faq/antoine-vapor-pressure.shtml What is the Antoine Equation?] (Chemistry Department, [[Frostburg State University]], [[Maryland]])</ref><ref name=Sinnot>{{cite book|author=R.K.Sinnot|title=[http://books.google.ca/books?id=DJaxUL3numgC&pg=PA331&lpg=PA331&dq=antoine+equation+constants&source=bl&ots=2c0cqzbR0t&sig=YPuvrW2kWnWP2s4QvY9TTpxGgNM&hl=en&sa=X&ei=widrUcu9Cce0qgHq24CgCA&ved=0CGcQ6AEwBzgK#v=onepage&q=antoine%20equation%20constants&f=false Chemical Engineering Design]|edition=4th|publisher=Butterworth-Heinemann|year=2005|page=331|isbn=0-7506-6538-6}}</ref> is a mathematical expression of the relation between the vapor pressure and the temperature of pure liquid or solid substances. The basic form of the equation is:
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| :<math>\log P = A-\frac{B}{C+T}</math>
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| and it can be transformed into this temperature-explicit form: | |
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| :<math>T = \frac{B}{A-\log P} - C</math>
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| where: <span style="vertical-align:+12%;"><math>P</math></span> is the absolute vapor pressure of a substance<br>
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| : <span style="vertical-align:+15%;"><span style="vertical-align:+12%;"><math>T</math></span> is the temperature of the substance</span>
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| : <span style="vertical-align:+12%;"><math>A</math></span>, <span style="vertical-align:+12%;"><math>B</math></span> and <span style="vertical-align:+12%;"><math>C</math></span> are substance-specific coefficients (i.e., constants or parameters)
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| : <span style="vertical-align:-30%;"><math>\log</math> is typically either <math>\log_{10}</math> or <math>\log_e</math></span><ref name=Sinnot/>
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| A simpler form of the equation with only two coefficients is sometimes used:
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| :<math>\log P = A-\frac{B}{T}</math>
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| which can be transformed to:
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| :<math>T = \frac{B}{A-\log P}</math>
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| Sublimations and vaporizations of the same substance have separate sets of Antoine coefficients, as do components in mixtures.<ref name=frostburg/> Each parameter set for a specific compound is only applicable over a specified temperature range. Generally, temperature ranges are chosen to maintain the equation's accuracy of a few up to 8-10 percent. For many volatile substances, several different sets of parameters are available and used for different temperature ranges. The Antoine equation has poor accuracy with any single parameter set when used from a compound's melting point to its critical temperature. Accuracy is also usually poor when vapor pressure is under 10 Torr because of the limitations of the apparatus used to establish the Antoine parameter values.
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| The Wagner Equation<ref>{{Citation|last= Wagner|first= W.|title= New vapour pressure measurements for argon and nitrogen and an new method for establishing rational vapour pressure equations|journal= Cryogenics|volume= 13|issue= 8|pages= 470–482 |year= 1973|doi= }}</ref> gives "one of the best"<ref>Perry's Chemical Engineers' Handbook, 7th Ed. pg 4-15</ref> fits to experimental data but is quite complex. It expresses reduced vapor pressure as a function of reduced temperature.
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| ==Relation to boiling point of liquids==
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| {{further|Boiling point}}
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| [[Image:Vapor Pressure Chart.png|thumb|right|301 px|A typical vapor pressure chart for various liquids]]
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| 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.<ref>{{cite book|author=Perry, R.H. and Green, D.W. (Editors)|title=[[Perry's Chemical Engineers' Handbook]]|edition=7th|publisher=McGraw-Hill|year=1997|isbn= 0-07-049841-5}}</ref>
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| 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 (−24.2 °C), which is where the vapor pressure curve of methyl chloride (the blue line) intersects the horizontal pressure line of one atmosphere ([[Atmosphere (unit)|atm]]) of absolute vapor pressure.
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| 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)<ref>{{cite article|author=Dreisbach, R. R. and Spencer, R. S.| title=Infinite Points of Cox Chart Families and dt/dP Values at any Pressure|journal=Industrial and Engineering Chemistry,|volume=41|number=1|page=176|month=January|year=1949}}</ref> 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.
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| ==Liquid mixtures==
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| [[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:
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| :<math> p_\text{tot} = \sum_i p_i \chi_i \,</math>
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| where '''''p'''''<sub> '''tot'''</sub> is the mixture's vapor pressure, '''''i''''' is one of the components of the mixture and '''''Χ<sub>i</sub>''''' is the [[mole fraction]] of that component in the liquid mixture. The term '''''p<sub>i</sub>Χ<sub>i</sub>''''' is the partial 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 force]]s).
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| 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.
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| 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.
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| The negative and positive deviations can be used to determine [[thermodynamic activity]] coefficients of the components of mixtures.
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| ==Solids==
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| [[Image:Vapor Pressure Curve of Liquid and Solid Benzene.png|thumb|300px|Vapor pressure of liquid and solid benzene]]
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| 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 (physics)|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, 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]].
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| 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 [[Enthalpy of fusion|heat of fusion]] is known, by using this particular form of the [[Clausius–Clapeyron relation]]:<ref name="Moller">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</ref>
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| :<math>\ln\,P^S_{solid} = \ln\,P^S_{liquid} - \frac{\Delta H_m}{R} \left( \frac{1}{T} - \frac{1}{T_m} \right)</math>
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| with:
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| {| border="0" cellpadding="1"
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| |-
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| !align=right|<math>P^S_{solid}</math>
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| |align=left|= Sublimation pressure of the solid component at the temperature <math>T<T_m</math>
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| |-
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| !align=right|<math>P^S_{liquid}</math>
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| |align=left|= Extrapolated vapor pressure of the liquid component at the temperature <math>T<T_m</math>
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| |-
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| !align=right|<math>\Delta H_m</math>
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| |align=left|= Heat of fusion
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| |-
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| !align=right|<math>R</math>
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| |align=left|= [[Gas constant]]
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| |-
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| !align=right|<math>T</math>
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| |align=left|= Sublimation temperature
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| |-
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| !align=right|<math>T_m</math>
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| |align=left|= Melting point temperature
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| |}
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| 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 (Δ''H''<sub>m</sub> is positive) and the difference grows with increased distance from the melting point.
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| ==Boiling point of water==
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| [[Image:Water vapor pressure graph.jpg|thumb|right|Graph of water vapor pressure versus temperature. At the normal boiling point of 100°C, it equals the standard atmospheric pressure of 760 Torr or 101.325 kPa.]]
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| {{main|Vapor pressure of water}}
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| 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]]:
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| :<math>\log_{10}P = 8.07131 - \frac{1730.63}{233.426 + T_b}</math>
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| or transformed into this temperature-explicit form:
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| :<math>T_b = \frac{1730.63}{8.07131 - \log_{10}P} - 233.426</math>
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| where the temperature <math>T_b</math> is the boiling point in degrees [[Celsius]] and the pressure <math>P_{ } </math> is in [[Torr]].
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| ==Dühring's rule==
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| {{main|Dühring's rule}}
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| Dühring's rule states that a linear relationship exists between the temperatures at which two solutions exert the same vapor pressure.
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| ==Examples==
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| The following table is a list of a variety of substances ordered by increasing vapor pressure.
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| {|style="text-align:center;" class="wikitable"
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| ! Substance
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| ! Vapor Pressure<br>(SI units)
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| ! Vapor Pressure<br>(Bar);
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| ! Vapor Pressure<br>(mmHg);
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| ! Temperature
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| |-
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| | [[Tungsten]]
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| | 100 Pa
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| | 0.001
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| | 0.75
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| | 3203 °C
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| |-
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| | [[Ethylene glycol]]
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| | 500 Pa
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| | 0.005
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| | 3.75
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| | 20 °C
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| |-
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| | [[Xenon difluoride]]
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| | 600 Pa
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| | 0.006
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| | 4.50
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| | 25 °C
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| |-
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| | [[Water]] (H<sub>2</sub>O)
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| | 2.3 kPa
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| | 0.023
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| | 17.5
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| | 20 °C
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| |-
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| | [[Propanol]]
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| | 2.4 kPa
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| | 0.024
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| | 18.0
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| | 20 °C
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| |-
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| | [[Ethanol]]
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| | 5.83 kPa
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| | 0.0583
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| | 43.7
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| | 20 °C
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| |-
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| | [[Methyl isobutyl ketone]]
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| | 2.66 kPa
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| | 0.0266
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| | 19.95
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| | 25 °C
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| |-
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| | [[Freon|Freon 113]]
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| | 37.9 kPa
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| | 0.379
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| | 284
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| | 20 °C
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| |-
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| | [[Acetaldehyde]]
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| | 98.7 kPa
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| | 0.987
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| | 740
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| | 20 °C
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| |-
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| | [[Butane]]
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| | 220 kPa
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| | 2.2
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| | 1650
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| | 20 °C
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| |-
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| | [[Formaldehyde]]
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| | 435.7 kPa
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| | 4.357
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| | 3268
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| | 20 °C
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| |-
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| | [[Propane]]
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| | 1.013 MPa
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| | 10.133
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| | 7600
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| | 25.6 °C
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| |-
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| | [[Carbonyl sulfide]]
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| | 1.255 MPa
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| | 12.55
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| | 9412
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| | 25 °C
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| |-
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| | [[Carbon dioxide]]
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| | 5.7 MPa
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| | 57
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| | 42753
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| | 20 °C
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| |-
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| |}
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| ==Estimating vapor pressure from molecular structure==
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| Several empirical methods exist to estimate liquid vapor pressure from molecular structure for organic molecules. Some examples are SIMPOL,<ref>{{cite journal|author=J. F. Pankow et al.|title=SIMPOL.1: a simple group contribution method for predicting vapor pressures and enthalpies of vaporization of multifunctional organic compounds|journal=Atmos. Chem. Phys.|volume=8|pages=2773–2796|year=2008|url=http://www.atmos-chem-phys.net/8/2773/2008/acp-8-2773-2008.html}}</ref> the method of Moller et al.,<ref name = "Moller" /> and EVAPORATION.<ref>[http://tropo.aeronomie.be/models/evaporation_run.htm "Vapour pressure of pure liquid compounds. Estimation by EVAPORATION"]</ref><ref>{{cite journal|author=S. Compernolle et al.|title=EVAPORATION: a new vapour pressure estimation method for organic molecules including non-additivity and intramolecular interactions|journal=Atmos. Chem. Phys.|volume=11|pages=9431–9450|year=2011|url=http://www.atmos-chem-phys.net/11/9431/2011/acp-11-9431-2011.html}}</ref>
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| ==Meaning in meteorology==
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| In [[meteorology]], the term ''vapor pressure'' is used to mean the [[partial pressure]] of [[water vapor]] in the atmosphere, even if it is not in equilibrium,<ref>[http://amsglossary.allenpress.com/glossary/search?id=vapor-pressure1 Glossary] (Developed by the [[American Meteorological Society]])</ref> and the ''equilibrium vapor pressure'' is specified otherwise. 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.<ref>[http://fermi.jhuapl.edu/people/babin/vapor/index.html A Brief Tutorial] (An article about the definition of equilibrium vapor pressure)</ref>
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| ==See also==
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| * [[Absolute humidity]]
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| * [[Clausius-Clapeyron relation]]
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| * [[Partial pressure]]
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| * [[Relative humidity]]
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| * [[Relative volatility]]
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| * [[Raoult's law]]
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| * [[Saturation vapor density]]
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| * [[Triple point]]
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| * [[Vapor-liquid equilibrium]]
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| * [[Vapour pressure of water|Vapor pressure of water]]
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| * [[Volatility (chemistry)|Volatility]]
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| * [[Reid vapor pressure]]
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| * [[True vapor pressure]]
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| * [[Vapor pressures of the elements (data page)]]
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| ==References==
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| {{Reflist}}
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| ==External links==
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| *[http://www.engineersedge.com/fluid_flow/fluid_data.htm Fluid Characteristics Chart]
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| *[http://hyperphysics.phy-astr.gsu.edu/hbase/kinetic/vappre.html#c2 Hyperphysics]
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| *[http://www.ilpi.com/msds/ref/vaporpressure.html MSDS Vapor Pressure]
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| *[http://www.envmodels.com/freetools.php?menu=pression Online vapor pressure calculation tool (Requires Registration)]
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| *[http://www.aim.env.uea.ac.uk/aim/ddbst/pcalc_main.php Prediction of Vapor Pressures of Pure Liquid Organic Compounds]
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| [[Category:Thermodynamic properties]]
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| [[Category:Chemical engineering]]
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| [[Category:Meteorology]]
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| [[Category:Gases]]
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| [[Category:Pressure]]
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| [[fr:Pression de vapeur saturante]]
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