Boiling point
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{{Otheruses1|the boiling point of liquids}}
The '''boiling point''' of a liquid is the temperature at which the [[vapor pressure]] of the liquid equals the environmental pressure surrounding the liquid.<ref>{{cite book|author=David.E. Goldberg|title=3,000 Solved Problems in Chemistry|edition=First Edition|publisher=McGraw-Hill|year=1988|id=ISBN 0-07-023684-4}} Section 17.43, page 321</ref><ref>{{cite book|author=Louis Theodore, R. Ryan Dupont and Kumar Ganesan (Editors)|title=Pollution Prevention: The Waste Management Approach to the 21st Century|edition= |publisher=CRC Press|year=1999|id=ISBN 1-56670-495-2}} Section 27, page 15</ref><ref>[http://www.airproducts.com/Products/fastfacts/charts_n_tables/definitions.asp Gas Products Charts and Tables - Definitions]</ref><ref>[http://lorien.ncl.ac.uk/ming/distil/distilpri.htm Distillation Principals] Ming T. Tham, Senior Lecturer, [[University of Newcastle upon Tyne]]</ref> A liquid in a vacuum environment has a lower boiling point than when the liquid is at [[atmospheric pressure]]. A liquid in a high pressure environment has a higher boiling point than when the liquid is at atmospheric pressure. In other words, the boiling point of liquids varies with and depends upon the surrounding environmental pressure.
The '''normal boiling point''' (also called the '''atmospheric boiling point''' or the '''atmospheric pressure boiling point''') of a liquid is the special case in which the vapor pressure of the liquid equals the defined atmospheric pressure at sea level, 1 atmosphere.<ref>[http://www.chem.purdue.edu/gchelp/gloss/normalbp.html General Chemistry Glossary] [[Purdue University]] website page</ref><ref>{{cite book|author=Kevin R. Reel, R. M. Fikar, P. E. Dumas, Jay M. Templin, and Patricia Van Arnum|title=AP Chemistry (REA) - The Best Test Prep for the Advanced Placement Exam|edition=9th Edition|publisher=Research & Education Association|year=2006|id=ISBN 0-7386-0221-3}} Section 71, page 224</ref> At that temperature, the vapor pressure of the liquid becomes sufficient to overcome atmospheric pressure and lift the liquid to form bubbles inside the bulk of the liquid. The '''standard boiling point''' is now (as of 1982) defined by [[IUPAC]] as the temperature at which boiling occurs under a pressure of 1 bar.<ref>[http://www.iupac.org/publications/pac/1982/pdf/5406x1239.pdf Notation for States and Processes, Significance of the Word Standard in Chemical Thermodynamics, and Remarks on Commonly Tabulated Forms of Thermodynamic Functions] See page 1274</ref>
The [[heat of vaporization]] is the amount of heat required to convert or vaporize a saturated liquid (i.e., a liquid at its boiling point) into a vapor.
Liquids may change to a vapor at temperatures below their boiling points through the process of [[evaporation]]. Evaporation is a surface phenomenon in which molecules located near the vapor/liquid surface escape into the vapor phase. On the other hand, boiling is a process in which molecules anywhere in the liquid escape, resulting in the formation of vapor bubbles within the liquid.
== Saturation temperature and pressure==
A '''saturated liquid''' contains as much thermal energy as it can without boiling (or conversely a '''saturated vapor''' contains as little thermal energy as it can without [[Condensation|condensing]]).
'''Saturation temperature''' means ''boiling point''. The saturation temperature is the temperature for a corresponding saturation pressure at which a liquid boils into its vapor [[phase (matter)|phase]]. The liquid can be said to be saturated with [[thermal energy]]. Any addition of thermal energy results in a [[phase transition]].
If the [[pressure]] in a system remains constant ([[Isobaric process|isobaric]]), a vapor at saturation temperature will begin to condense into its liquid phase as thermal energy ([[heat]]) is removed. Similarly, a liquid at saturation temperature and pressure will boil into its vapor phase as additional thermal energy is applied.
The boiling point corresponds to the temperature at which the vapor pressure of the liquid equals the surrounding environmental pressure. Thus, the boiling point is dependent on the pressure. Usually, boiling points are published with respect to atmospheric pressure (101.325 [[kilopascal]]s or 1 [[atmospheric pressure|atm]]). At higher elevations, where the atmospheric pressure is much lower, the boiling point is also lower. The boiling point increases with increased pressure up to the [[critical point (thermodynamics)|critical point]], where the gas and liquid properties become identical. The boiling point cannot be increased beyond the critical point. Likewise, the boiling point decreases with decreasing pressure until the [[triple point]] is reached. The boiling point cannot be reduced below the triple point.
If the heat of vaporization and the vapor pressure of a liquid at a certain temperature is known, the normal boiling point can be calculated by using the [[Clausius-Clapeyron equation]] thus:
<math>T_B = \Bigg(\frac{\,R\,[\,\ln(P_0)-\ln(101.325)\,]}{\Delta H_{vap}}+\frac{1}{T_0}\Bigg)^{-1}</math>
{| border="0" cellpadding="2"
|-
|align=right|where:
|
|-
!align=right|<math>T_B</math>
|align=left|= the normal boiling point, K
|-
!align=right|<math>R</math>
|align=left|= the [[ideal gas constant]], 8.314 J · K<sup>-1</sup> · mol<sup>-1</sup>
|-
!align=right|<math>P_0</math>
|align=left|= is the vapor pressure at a given temperature, kPa
|-
!align=right| <math>101.325</math>
|align=left|= atmospheric pressure, kPa
|-
!align=right|<math>\Delta H_{vap} </math>
|align=left|= the heat of vaporization of the liquid, J/mol
|-
!align=right|<math>T_0</math>
|align=left|= the given temperature, K
|-
!align=right|<math>ln</math>
|align=left|= the [[natural logarithm]] to the base [[e]]
|}
'''Saturation pressure''' is the pressure for a corresponding saturation temperature at which a liquid boils into its vapor phase. Saturation pressure and saturation temperature have a direct relationship: as saturation pressure is increased so is saturation temperature.
If the temperature in a [[system]] remains constant (an ''[[isothermal]]'' system), vapor at saturation pressure and temperature will begin to [[condensation|condense]] into its liquid phase as the system pressure is increased. Similarly, a liquid at saturation pressure and temperature will tend to [[Flash evaporation|flash]] into its vapor phase as system pressure is decreased.
The boiling point of [[water]] is 100 [[Celsius|°C]] (212 [[Fahrenheit|°F]]) at standard pressure. On top of [[Mount Everest]] the pressure is about 260 [[Millibar|mbar]] (26.39 kPa) so the boiling point of water is 69 [[Celsius|°C]]. (156.2 [[Fahrenheit|°F]]).
For purists, the ''normal boiling point of water'' is 99.97 degrees Celsius at a pressure of 1 atm (i.e., 101.325 kPa). Until 1982 this was also the ''standard boiling point of water'', but the [[IUPAC]] now recommends a standard pressure of 1 bar (100 kPa). At this slightly reduced pressure, the ''standard boiling point of water'' is 99.61 degrees Celsius.
==Relation between the normal boiling point and the vapor pressure of liquids ==
[[Image:Vapor Pressure Chart.png|thumb|right|301 px|A typical vapor pressure chart for various 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.<ref>{{cite book|author=Perry, R.H. and Green, D.W. (Editors)|title=[[Perry's Chemical Engineers' Handbook]]|edition=7th Edition|publisher=McGraw-Hill|year=1997|id= ISBN 0-07-049841-5}}</ref> 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(-43.7 °C), which is where the vapor pressure curve of propane (the purple line) intersects the horizontal pressure line of one atmosphere ([[Atmosphere (unit)|atm]]) of absolute vapor pressure.
In terms of [[intermolecular]] interactions, the boiling point represents the point at which the liquid [[molecules]] possess enough [[thermal energy]] to overcome the various intermolecular attractions binding the molecules into the liquid (eg. [[dipole-dipole attraction]], [[instantaneous-dipole induced-dipole attraction]]s, and [[hydrogen bond]]s). Therefore the boiling point is also an indicator of the strength of these attractive forces.
== Properties of the elements ==
{{see|List of elements by boiling point}}
The element with the lowest boiling point is [[helium]]. Both the boiling points of [[rhenium]] and [[tungsten]] exceed 5000 [[kelvin|K]] at [[standard pressure]]. Due to the experimental difficulty of precisely measuring extreme temperatures without bias, there is some discrepancy in the literature as to whether [[tungsten]] or [[rhenium]] has the higher boiling point.<ref>{{cite book|author=Howard DeVoe|title=Thermodynamics and Chemistry|edition=1st Edition|publisher=Prentice-Hall|date=2000|id=ISBN 0-02-328741-1}}</ref>
==See also==
* [[Boiling delay]]
* [[Boiling-point elevation]]
* [[Critical temperature]]
* [[Joback method]] (Estimation of normal boiling points from molecular structure)
==External Links==
*[http://www2.sigmaaldrich.com/suite7/Area_of_Interest/Research_Essentials/Solvents/Key_Resources/nomograph.html?cm_mmc=wiki-_-social-_-nomograph-_-boiling_point Sigma-Aldrich Pressure-Temperature Nomograph] Quickly and easily estimate boiling points at various pressures.
==References==
{{reflist}}
{{Phase of matter}}
[[Category:Thermodynamics]]
[[Category:Fundamental physics concepts]]
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