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{{About|the boiling point of liquids}}
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[[Image:Kochendes wasser02.jpg|thumb|300px|Boiling water]]The '''boiling point''' of a substance is the temperature at which the [[vapor pressure]] of the [[liquid]] equals the pressure surrounding the liquid<ref>{{cite book|author=David.E. Goldberg|title=3,000 Solved Problems in Chemistry|edition=1st|publisher=McGraw-Hill|year=1988|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|publisher=CRC Press|year=1999|isbn=1-56670-495-2}} Section 27, page 15</ref> and the liquid changes into a vapor.
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A liquid in a [[vacuum]] has a lower boiling point than when that liquid is at [[atmospheric pressure]]. A liquid at high-[[pressure]] has a higher boiling point than when that liquid is at atmospheric pressure. In other words, the boiling point of a liquid varies depending upon the surrounding environmental pressure. For a given pressure, different liquids boil at different temperatures.
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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 (unit)|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|publisher=Research & Education Association|year=2006|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 allow bubbles of vapor to form 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 (unit)|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>
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The [[heat of vaporization]] is the energy required to transform a given quantity (a mol, kg, pound, etc.) of a substance from a liquid into a gas at a given pressure (often atmospheric pressure).
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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 liquid's edge, not contained by enough liquid pressure on that side, escape into the surroundings as [[vapor]]. 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.
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==Saturation temperature and pressure==
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A ''[[Saturation (chemistry)|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]]).
 
 
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'''Saturation temperature''' means ''boiling point''. The saturation temperature is the temperature for a corresponding saturation pressure at which a liquid boils into its [[Gas|vapor phase]]. The liquid can be said to be saturated with [[thermal energy]]. Any addition of thermal energy results in a [[phase transition]].
 
 
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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.
 
 
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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.
 
 
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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:
 
 
</ul>
<math>T_B = \Bigg(\frac{\,R\,\ln(P_0)}{\Delta H_{vap}}+\frac{1}{T_0}\Bigg)^{-1}</math>
 
{| border="0" cellpadding="2"
|-
|align=right|where:
|&nbsp;
|-
!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, atm
|-
!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 (mathematical constant)|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 [[Properties of water|water]] is 100&nbsp;[[Celsius|°C]] (212&nbsp;[[Fahrenheit|°F]]) at standard pressure. On top of [[Mount Everest]], at {{convert|8848|m|ft|abbr=on}} elevation, the pressure is about {{convert|252|torr|kPa|3|lk=on|abbr=on}}<ref>[http://jap.physiology.org/content/86/3/1062.long Barometric pressure on Mt Everest: new data and physiological significance]</ref> and the boiling point of water is {{convert|71|°C|°F|1|lk=on}}. The boiling point decreases 1 °C every 285 m of elevation, or 1 °F every 500&nbsp;ft.
 
There are two conventions regarding the ''standard boiling point of water'': The normal boiling point 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).<ref>[http://goldbook.iupac.org/S05921.html Standard Pressure] IUPAC defines the "standard pressure" as being 10<sup>5</sup> Pa (which amounts to 1 bar).</ref> At this slightly reduced pressure, the ''standard boiling point of water'' is 99.61 degrees Celsius.<ref>[http://highered.mcgraw-hill.com/sites/dl/free/0073529214/395307/appdxs1_2.pdf Appendix 1: Property Tables and Charts (SI Units)], Scroll down to Table A-5 and read the temperature value of 99.61 °C at a pressure of 100 kPa (1 bar). Obtained from McGraw-Hill's Higher Education website.</ref>
 
==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|publisher=McGraw-Hill|year=1997|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, [[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.
 
== Properties of the elements ==
{{further|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]]; because it is difficult to measure extreme temperatures precisely without bias, both have been cited in the literature as having the higher boiling point.<ref>{{cite book|author=Howard DeVoe|title=Thermodynamics and Chemistry|edition=1st|publisher=Prentice-Hall|year=2000|isbn=0-02-328741-1}}</ref>
 
==Boiling point as a reference property of a pure compound==
As can be seen from the above plot of the logarithm of the vapor pressure vs. the temperature for any given pure [[chemical compound]], its normal boiling point can serve as an indication of that compound's overall [[Volatility (chemistry)|volatility]]. A given pure compound has only one normal boiling point, if any, and a compound's normal boiling point and [[melting point]] can serve as characteristic [[Physical property|physical properties]] for that compound, listed in reference books. The higher a compound's normal boiling point, the less volatile that compound is overall, and conversely, the lower a compound's normal boiling point, the more volatile that compound is overall.  Some compounds decompose at higher temperatures before reaching their normal boiling point, or sometimes even their melting point. For a stable compound, the boiling point ranges from its [[triple point]] to its [[Critical point (thermodynamics)|critical point]], depending on the external pressure.  Beyond its triple point, a compound's normal boiling point, if any, is higher than its melting point.  Beyond the critical point, a compound's liquid and vapor phases merge into one phase, which may be called a superheated gas.  At any given temperature, if a compound's normal boiling point is lower, then that compound will generally exist as a gas at atmospheric external pressure.  If the compound's normal boiling point is higher, then that compound can exist as a liquid or solid at that given temperature at atmospheric external pressure, and will so exist in equilibrium with its vapor (if volatile) if its vapors are contained.   If a compound's vapors are not contained, then some volatile compounds can eventually evaporate away in spite of their higher boiling points.
 
In general, compounds with [[ionic bond]]s have high normal boiling points, if they do not decompose before reaching such high temperatures. Many [[metal]]s have high boiling points, but not all. Very generally&mdash;with other factors being equal&mdash;in compounds with covalently bonded [[molecule]]s, as the size of the molecule (or [[molecular mass]]) increases, the normal boiling point increases. When the molecular size becomes that of a [[macromolecule]], [[polymer]], or otherwise very large, the compound often decomposes at high temperature before the boiling point is reached.  Another factor that affects the normal boiling point of a compound is the [[Polarity (chemistry)|polarity]] of its molecules. As the polarity of a compound's molecules increases, its normal boiling point increases, other factors being equal. Closely related is the ability of a molecule to form [[hydrogen bond]]s (in the liquid state), which makes it harder for molecules to leave the liquid state and thus increases the normal boiling point of the compound. Simple [[carboxylic acid]]s dimerize by forming hydrogen bonds between molecules. A minor factor affecting boiling points is the shape of a molecule. Making the shape of a molecule more compact tends to lower the normal boiling point slightly compared to an equivalent molecule with more surface area.
 
[[File:Binary Boiling Point Diagram new.svg|thumb|right|350px|'''Binary boiling point diagram''' of two hypothetical only weakly interacting components without an azeotrope]]
Most volatile compounds (anywhere near ambient temperatures) go through an intermediate liquid phase while warming up from a solid phase to eventually transform to a vapor phase. By comparison to boiling, a [[Sublimation (phase transition)|sublimation]] is a physical transformation in which a solid turns directly into vapor, which happens in a few select cases such as with [[carbon dioxide]] at atmospheric pressure. For such compounds, a [[sublimation point]] is a temperature at which a solid turning directly into vapor has a vapor pressure equal to the external pressure.
 
==Impurities and mixtures==
In the preceding section, boiling points of pure compounds were covered. Vapor pressures and boiling points of substances can be affected by the presence of dissolved impurities ([[Solution|solutes]]) or other miscible compounds, the degree of effect depending on the concentration of the impurities or other compounds. The presence of non-volatile impurities such as [[Salt (chemistry)|salts]] or compounds of a [[Volatility (chemistry)|volatility]] far lower than the main component compound decreases its [[mole fraction]] and the [[Solution (chemistry)|solution's]] volatility, and thus raises the normal boiling point in proportion to the [[concentration]] of the solutes. This effect is called '''[[boiling point elevation]]'''. As a common example, [[Saline water|salt water]] boils at a higher temperature than pure [[water]].
 
In other mixtures of miscible compounds (components), there may be two or more components of varying volatility, each having its own pure component boiling point at any given pressure.  The presence of other volatile components in a mixture affects the vapor pressures and thus boiling points and [[dew point]]s of all the components in the mixture. The dew point is a temperature at which a vapor [[Condensation|condenses]] into a liquid.   Furthermore, at any given temperature, the composition of the vapor is different from the composition of the liquid in most such cases. In order to illustrate these effects between the volatile components in a mixture, a '''boiling point diagram''' is commonly used. [[Distillation]] is a process of boiling and [usually] condensation which takes advantage of these differences in composition between liquid and vapor phases.
 
==See also==
* [[Boiling points of the elements (data page)]]
* [[Boiling-point elevation]]
* [[Critical point (thermodynamics)]]
* [[Ebulliometer]]
* [[Joback method]] (Estimation of normal boiling points from molecular structure)
* [[Subcooling]]
* [[Superheating]]
* [[Trouton's constant]]
 
==References==
{{Reflist}}
 
{{Phase of matter}}
 
{{DEFAULTSORT:Boiling Point}}
[[Category:Concepts in physics]]
[[Category:Thermodynamics]]
[[Category:Threshold temperatures]]
 
[[el:Σημείο βρασμού]]

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Okay so cancer shield aside Air Jordan Retro

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Okay so cancer shield aside, let's discuss Armstrong's fall from grace. Similar to Tiger Wood's tumult, although for vastly different reasons, once more a national hero becomes tabloid fodder. Armstrong's scandal has been described as "the most sophisticated, professionalized, and successful doping program." I suppose anything worth doing is Air Jordan Retro worth doing well, but I don't really care that he cheated. I think it's wrong to cheat however i am not surprised when professional athletes are subjected for using any means possible once the institution of sport almost endorses it.

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