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| {{About|one of the gas laws in physics|the chemist|William Henry (chemist)|all other uses|Law (disambiguation)}}
| | I am 20 years old and my name is Lucie Biaggini. I life in Freudenthal (Austria).<br><br>Also visit my blog post - [http://hemorrhoidtreatmentfix.com/internal-hemorrhoids-treatment internal hemorrhoid treatment] |
| In [[physics]], '''Henry's law''' is one of the [[gas laws]] formulated by [[William Henry (chemist)|William Henry]] in 1803. It states:
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| :"At a constant temperature, the amount of a given gas that dissolves in a given type and volume of liquid is directly proportional to the [[partial pressure]] of that gas in equilibrium with that liquid."
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| An equivalent way of stating the law is that ''the [[solubility]] of a gas in a liquid is directly proportional to the partial pressure of the gas above the liquid.''
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| An everyday example of Henry's law is given by [[Carbonation|carbonated]] [[soft drink]]s. Before the bottle or can of carbonated drink is opened, the gas above the drink is almost pure [[carbon dioxide]] at a pressure slightly higher than [[atmospheric pressure]]. The drink itself contains dissolved carbon dioxide. When the bottle or can is opened, some of this gas escapes, giving the characteristic hiss. Because the partial pressure of carbon dioxide above the liquid is now lower, some of the dissolved carbon dioxide comes out of solution as bubbles. If a glass of the drink is left in the open, the concentration of carbon dioxide in solution will come into equilibrium with the carbon dioxide in the air, and the drink will go "flat".
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| A slightly more exotic example of Henry's law is in the [[Decompression (diving)|decompression]] and [[decompression sickness]] of [[Underwater diving|underwater divers]].
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| ==Formula and the Henry's law constant==
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| Henry's law can be put into mathematical terms (at constant temperature) as
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| :<math>p = k_{\mathrm{H}} c</math>
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| where ''p'' is the [[partial pressure]] of the [[Solution|solute]] in the gas above the solution, ''c'' is the [[concentration]] of the solute and ''k''<sub>H</sub> is a constant with the [[Dimensional analysis|dimensions]] of pressure divided by concentration. The constant, known as the Henry's law constant, depends on the solute, the solvent and the temperature.
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| Some values for ''k''<sub>H</sub> for gases dissolved in [[water (molecule)|water]] at 298 [[kelvin|K]] include:
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| :[[oxygen]] (O<sub>2</sub>) : 769.2 [[Litre|L]]·[[Atmosphere (unit)|atm]]/[[Mole (unit)|mol]]
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| :[[carbon dioxide]] (CO<sub>2</sub>) : 29.41 L·atm/mol
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| :[[hydrogen]] (H<sub>2</sub>) : 1282.1 L·atm/mol
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| There are various other forms of Henry's Law which define the constant ''k''<sub>H</sub> differently and require different dimensional units.<ref name=SmithandHarvey>{{cite journal|author=Francis L. Smith and Allan H. Harvey |date=September 2007 |title=[http://www.chemengr.ucsb.edu/~ceweb/courses/che128/pdf/090733%20Avoid%20Common%20Pitfalls.pdf Avoid Common Pitfalls When Using Henry's Law]|journal=CEP (Chemical Engineering Progress) |volume= |issue= |pages= |issn=0360-7275}}</ref> In particular, the "concentration" of the solute in solution may also be expressed as a [[mole fraction]] or as a [[molarity]].<ref name=Lee>F.F. Lee (2007). [http://books.google.com/books?id=EHYYy_MDXE0C&printsec=frontcover#v=onepage&q&f=false ''Comprehensive analysis, Henry's law constant determination, and photocatalytic degradation of polychlorinated biphenyls (PCBs) and/or other persistent organic pollutants (POPs)''], Ph.D. dissertation, State University of New York at Albany, pp. 199-201.</ref>
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| ===Other forms of Henry's law===
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| The various other forms of Henry's law are discussed in the technical literature.<ref name=SmithandHarvey/><ref name=NCSU>[http://wikis.lib.ncsu.edu/index.php/CH_431/Lecture_14#Henry.27s_law North Carolina State University CH 431/Lecture 14]</ref><ref name="Sander">[http://www.henrys-law.org An extensive list of Henry's law constants, and a conversion tool]</ref>
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| {| class="wikitable" width="525"
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| |+ '''Table 1: Some forms of Henry's law and constants (gases in water at 298.15 K)<ref name="Sander"/>
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| ! equation: || <math>k_{\mathrm{H,pc}} = \frac{p}{c_\mathrm{aq}}</math> || <math>k_{\mathrm{H,cp}} = \frac{c_\mathrm{aq}}{p}</math> || <math> k_{\mathrm{H,px}} = \frac{p}{x} </math> || <math> k_{\mathrm{H,cc}} = \frac{c_{\mathrm{aq}}}{c_{\mathrm{gas}}} </math>
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| |-
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| ! units: || <math>\frac{\mathrm{L} \cdot \mathrm{atm}}{\mathrm{mol}}</math> || <math> \frac{\mathrm{mol}}{\mathrm{L} \cdot \mathrm{atm}}</math> || <math>\rm atm\,</math> || [[Dimensionless quantity|dimensionless]]
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| |-
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| |align=center| [[Oxygen|O<sub>2</sub>]] ||align=center| 769.23||align=center| 1.3{{e|−3}} ||align=center| 4.259{{e|4}} ||align=center| 3.181{{e|−2}}
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| |-
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| |align=center| [[Hydrogen|H<sub>2</sub>]] ||align=center| 1282.05 ||align=center| 7.8{{e|−4}} ||align=center| 7.099{{e|4}} ||align=center| 1.907{{e|−2}}
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| |-
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| |align=center| [[Carbon dioxide|CO<sub>2</sub>]] ||align=center| 29.41 ||align=center| 3.4{{e|−2}} ||align=center| 0.163{{e|4}} ||align=center| 0.8317
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| |-
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| |align=center| [[Nitrogen|N<sub>2</sub>]] ||align=center| 1639.34 ||align=center| 6.1{{e|−4}} ||align=center| 9.077{{e|4}} ||align=center| 1.492{{e|−2}}
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| |-
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| |align=center| [[Helium|He]] ||align=center| 2702.7 ||align=center| 3.7{{e|−4}}||align=center| 14.97{{e|4}} ||align=center| 9.051{{e|−3}}
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| |-
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| |align=center| [[Neon|Ne]] ||align=center| 2222.22 ||align=center| 4.5{{e|−4}} ||align=center| 12.30{{e|4}} ||align=center| 1.101{{e|−2}}
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| |-
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| |align=center| [[Argon|Ar]] ||align=center| 714.28 ||align=center| 1.4{{e|−3}} ||align=center| 3.955{{e|4}} ||align=center| 3.425{{e|−2}}
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| |-
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| |align=center| [[Carbon monoxide|CO]] ||align=center| 1052.63 ||align=center| 9.5{{e|−4}} ||align=center| 5.828{{e|4}} ||align=center| 2.324{{e|−2}}
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| |}
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| {|border="0" cellpadding="2"
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| |-
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| |align=right|where:
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| |align=right|'''''c'''''<sub>'''aq'''</sub>
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| |align=left|= [[molar concentration|concentration]] (or [[molarity]]) of gas in solution (in mol/L)
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| |-
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| |align=right|'''''c'''''<sub>'''gas'''</sub>
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| |align=left|= concentration of gas above the solution (in mol/L)
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| |-
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| |align=right|'''''p'''''
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| |align=left|= [[partial pressure]] of gas above the solution (in [[atmosphere (unit)|atm]])
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| |-
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| |align=right|'''''x'''''
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| |align=left|= [[mole fraction]] of gas in solution (dimensionless)
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| |}
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| As can be seen by comparing the equations in the above table, the Henry's law constant '''''k'''''<sub>'''H,pc'''</sub> is simply the inverse of the constant '''''k'''''<sub>'''H,cp'''</sub>. Since all '''''k'''''<sub>'''H</sub>''' may be referred to as Henry's law constants, readers of the technical literature must be quite careful to note which version of the equation is being used.<ref name=SmithandHarvey/>
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| It should also be noted, the Henry's law is a limiting law that only applies for 'sufficiently dilute' solutions. The range of concentrations in which it applies becomes narrower the more the system diverges from ideal behavior. Roughly speaking, that is the more chemically 'different' the solute is from the solvent. Typically, Henry's law is only applicable to gas solute mole fractions less than 0.03.<ref>{{cite book|author=Prausnitz et al| title=[http://books.google.com/books?id=VSwc1XUmYpcC&printsec=frontcover&dq=intitle:Molecular+intitle:Thermodynamics+intitle:of+intitle:Fluid+intitle:Phase+intitle:Equilibria&source=bl&ots=uYqPECCSfR&sig=R0tEs04XFLVTLMeyOA4RhmpH-84&hl=en&sa=X&ei=n51gULSFKojc2gXJiICoDA&ved=0CC8Q6AEwAA#v=onepage&q&f=false Molecular Thermodynamics of Fluid Phase Equilibria]|edition=3rd|publisher=Prentice Hall|year=1999|page=586|ISBN=0-13-977745-8}}</ref>
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| It also only applies simply for solutions where the solvent does not [[chemical reaction|react chemically]] with the gas being dissolved. A common example of a gas that does react with the solvent is [[carbon dioxide]], which forms [[carbonic acid]] (H<sub>2</sub>CO<sub>3</sub>) to a certain degree with water.
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| ===Temperature dependence of the Henry constant===
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| When the temperature of a system changes, the Henry constant will also change.<ref name=SmithandHarvey/> This is why some people prefer to name it Henry coefficient. Multiple equations assess the effect of temperature on the constant. These forms of the [[van 't Hoff equation]] are examples:<ref name="Sander"/>
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| :<math> k_{\rm H,pc}(T) = k_{\rm H,pc}(T^\ominus)\, \exp{ \left[ -C \, \left( \frac{1}{T}-\frac{1}{T^\ominus}\right)\right]}\, </math>
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| :<math> k_{\rm H,cp}(T) = k_{\rm H,cp}(T^\ominus)\, \exp{ \left[ C \, \left( \frac{1}{T}-\frac{1}{T^\ominus}\right)\right]}\, </math>
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| where
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| :'''''k'''''<sub>H</sub> for a given temperature is Henry's constant (as defined in this article's first section). Note that the sign of C depends on whether '''''k'''''<sub>H,pc</sub> or '''''k'''''<sub>H,cp</sub> is used.
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| :'''''T''''' is any given temperature, in K
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| :'''''T'''''<sup> <s>o</s></sup> refers to the standard temperature (298 K).
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| This equation is only an approximation, and should be used only when no better, experimentally derived formula is known for a given gas.
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| The following table lists some values for constant '''''C''''' (in Kelvins) in the equation above:
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| {|class="wikitable"
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| |+ '''Table 2: Values of ''C'' (in K)'''
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| | '''Gas''' || align=center| [[Oxygen|O<sub>2</sub>]] || align=center| [[Hydrogen|H<sub>2</sub>]] || align=center| [[CO2|CO<sub>2</sub>]] || align=center| [[Nitrogen|N<sub>2</sub>]] || align=center| [[Helium|He]] || align=center| [[Neon|Ne]] || align=center| [[Argon|Ar]] || align=center| [[Carbon monoxide|CO]]
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| |-
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| |align=center|'''''C''''' || align=center| 1700 ||align=center| 500 || align=center| 2400 || align=center| 1300 || align=center| 230 || align=center| 490 || align=center| 1300 || align=center| 1300
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| |}
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| Because solubility of permanent gases usually decreases with increasing temperature at around the room temperature, the partial pressure a given gas concentration has in liquid must increase. While heating water (saturated with nitrogen) from 25 to 95 °C, the solubility will decrease to about 43% of its initial value. This can be verified when heating water in a pot; small bubbles evolve and rise long before the water reaches boiling temperature. Similarly, carbon dioxide from a [[Carbonation|carbonated]] drink escapes much faster when the drink is not cooled because the required partial pressure of CO<sub>2</sub> to achieve the same solubility increases in higher temperatures. Partial pressure of CO<sub>2</sub> in the gas phase in equilibrium with seawater doubles with every 16 K increase in temperature.<ref>{{cite journal|doi=10.1016/S0967-0645(02)00003-6|pages= 1601– 1622|title=Global sea–air CO2 flux based on climatological surface ocean pCO2, and seasonal biological and temperature effects|year=2002|last1=Takahashi|first1=Taro|last2=Sutherland|first2=Stewart C.|last3=Sweeney|first3=Colm|last4=Poisson|first4=Alain|last5=Metzl|first5=Nicolas|last6=Tilbrook|first6=Bronte|last7=Bates|first7=Nicolas|last8=Wanninkhof|first8=Rik|last9=Feely|first9=Richard A.|journal=Deep Sea Research Part II: Topical Studies in Oceanography|volume=49|issue=9–10|displayauthors=9}}</ref>
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| The constant '''''C''''' may be regarded as:
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| :<math> C = -\frac{\Delta_{\rm solv}H}{R} = -\frac{{\rm d}\left[ \ln k_{\rm H}(T)\right]}{{\rm d}(1/T)}</math>
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| where
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| :'''Δ'''<sub>'''solv'''</sub>'''''H''''' is the [[enthalpy of solution]]
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| :'''''R''''' is the [[gas constant]].
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| The solubility of gases does not always decrease with increasing temperature. For aqueous solutions, the Henry's law constant usually goes through a maximum (i.e., the solubility goes through a minimum). For most permanent gases, the minimum is below 120 °C. Often, the smaller the gas molecule (and the lower the gas solubility in water), the lower the temperature of the maximum of the Henry's law constant. Thus, the maximum is about 30 °C for helium, 92 to 93 °C for argon, nitrogen and oxygen, and 114 °C for xenon.<ref>{{cite book|author=Editor: P. Cohen|title=[http://books.google.com/books?id=gqr8AAAACAAJ&dq=intitle:The+intitle:ASME+intitle:handbook+intitle:on+intitle:Water+intitle:Technology+intitle:for+intitle:Thermal+intitle:Power+intitle:Systems&source=bl&ots=4hLppJxbrT&sig=uR6uaS1EIQIRXiX-UUJKQDQ3Nms&hl=en&sa=X&ei=GrVgULHIH6jYywGr3oHwCw&ved=0CC8Q6AEwAA The ASME Handbook on Water Technology for Thermal Power Systems]|edition=|publisher=The American Society of Mechanical Engineers|year=1989|page=442|isbn=978-0-7918-0634-0 }}</ref>
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| ==Influence of electrolytes==
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| The influence of electrolytes on the solubility of gases is sometimes given by Sechenov (often spelled Setchenov) equation which accounts for the "[[salting out]]" (i.e., decreasing the solubility) or "[[salting in]]" (i.e., increasing the [[solubility]]) effect (see the article on [[Activity_coefficient#Measurement_and_prediction_of_activity_coefficients|activity coefficient]]). The Sechenov equation can be written as:<ref>{{cite book|author=Editor: Trevor M. Letcher|title=[http://books.google.ca/books?id=0VMd-M2KbuYC&pg=PA71&lpg=PA71&dq=Sechenov+equation&source=bl&ots=jdaw-sEMmZ&sig=JVGs3e8_1tnq3ay0ns6Be038XyE&hl=en&sa=X&ei=7oNeT7T7FqHw0gHry6i_Bw&redir_esc=y#v=onepage&q=Sechenov%20equation&f=false Developments and applications in solubility]|edition=1st|publisher=Royal Society of Chemistry|year=2007|page=71|ISBN=978-0854043729}}</ref>
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| ::<math>\log(*z_1/z_1) = k_{syz} y</math>
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| where:
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| * *z<sub>1</sub> is the solubility of gas 1 in pure solvent
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| * z<sub>1</sub> is the solubility of gas 1 in an electrolyte solution
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| * y expresses the salt composition
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| ==In geophysics==
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| In [[geophysics]], a version of Henry's law applies to the solubility of a [[noble gas]] in contact with [[silicate]] melt. One equation used is
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| :<math>C_{\rm melt}/C_{\rm gas} = \exp\left[-\beta(\mu^{\rm E}_{\rm melt} - \mu^{\rm E}_{\rm gas})\right]\,</math>
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| where:
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| :''C'' = the [[number concentration]]s of the solute gas in the melt and gas phases
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| :''β'' = 1/''k''<sub>B</sub>''T'', an inverse temperature scale: ''k''<sub>B</sub> = the [[Boltzmann constant]]
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| :''µ''<sup>E</sup> = the excess [[chemical potential]]s of the solute gas in the two phases.
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| ==Comparison to Raoult's law==
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| For a dilute solution, the concentration of the solute is approximately proportional to its [[mole fraction]] ''x'', and Henry's law can be written as:
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| :<math>p = k_{\rm H}\,x</math>
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| This can be compared with [[Raoult's law]]:
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| :<math>p = p^\star\,x</math>
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| where ''p''* is the vapor pressure of the pure component.
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| At first sight, Raoult's law appears to be a special case of Henry's law where ''k''<sub>H</sub> = ''p''*. This is true for pairs of closely related substances, such as [[benzene]] and [[toluene]], which obey Raoult's law over the entire composition range: such mixtures are called "ideal mixtures".
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| The general case is that both laws are [[Limit of a function|limit laws]], and they apply at opposite ends of the composition range. The vapor pressure of the component in large excess, such as the solvent for a dilute solution, is proportional to its mole fraction, and the constant of proportionality is the vapor pressure of the pure substance (Raoult's law). The vapor pressure of the solute is also proportional to the solute's mole fraction, but the constant of proportionality is different and must be determined experimentally (Henry's law). In mathematical terms:
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| :Raoult's law: <math>\lim_{x\to 1}\left( \frac{p}{x}\right) = p^\star</math>
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| :Henry's law: <math>\lim_{x\to 0}\left( \frac{p}{x}\right) = k_{\rm H}</math>
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| Raoult's law can also be related to non-gas solutes.
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| ==Standard chemical potential==
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| Henry's law has been shown to apply to a wide range of solutes in the limit of "infinite dilution" (''x''→0), including non-volatile substances such as [[sucrose]] or even [[sodium chloride]]. In these cases, it is necessary to state the law in terms of [[chemical potential]]s. For a solute in an ideal dilute solution, the chemical potential depends on the concentration:
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| :<math>\mu = \mu_c^\ominus + RT\ln{\left( \frac{\gamma_c c}{c^\ominus}\right)}\,</math>, where <math>\gamma_c = \frac{k_{{\rm H,}c}}{p^\star}</math> for a volatile solute; ''c''<sup><s>o</s></sup> = 1 mol/L.
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| For non-ideal solutions, the [[activity coefficient]] ''γ<sub>c</sub>'' depends on the concentration and must be determined at the concentration of interest. The activity coefficient can also be obtained for non-volatile solutes, where the vapor pressure of the pure substance is negligible, by using the [[Gibbs–Duhem relation]]:
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| :<math>\sum_i n_i\, {\rm d}\mu_i = 0</math>
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| By measuring the change in vapor pressure (and hence chemical potential) of the solvent, the chemical potential of the solute can be deduced.
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| The [[standard state]] for a dilute solution is also defined in terms of infinite-dilution behavior. Although the standard concentration ''c''<sup><s>o</s></sup> is taken to be 1 mol/l by convention, the standard state is a hypothetical solution of 1 mol/l in which the solute has its limiting infinite-dilution properties. This has the effect that all non-ideal behavior is described by the activity coefficient: the activity coefficient at 1 mol/l is not necessarily unity (and is frequently quite different from unity).
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| All the relations above can also be expressed in terms of [[molality|molalities]] ''b'' rather than concentrations, e.g.:
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| :<math>\mu = \mu_b^\ominus + RT\ln{\left( \frac{\gamma_b b}{b^\ominus}\right)}\,</math>, where <math>\gamma_b = \frac{k_{{\rm H,}b}}{p^\star}</math> for a volatile solute; ''b''<sup><s>o</s></sup> = 1 mol/kg.
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| The standard chemical potential ''μ<sub>m</sub>''<sup><s>o</s></sup>, the activity coefficient ''γ<sub>m</sub>'' and the Henry's law constant ''k''<sub>H,''b''</sub> all have different numerical values when molalities are used in place of concentrations.
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| ==See also==
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| *[[Bunsen solubility coefficient]]
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| *[[Dalton's law]]
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| *[[Partial pressure]]
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| *[[Pervaporation]]
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| *[[Sieverts' law]]
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| *[[Henry adsorption constant]]
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| ==References==
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| {{reflist}}
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| ==External links==
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| * [http://www.composite-agency.com/messages/3724.html Ethanol solubility in EPDM], Solubility of chemicals in polymers using Henry's law
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| * {{Cite journal|doi=10.1021/cr60242a003|title=The Solubility of Gases in Liquids|year=1966|last1=Battino|first1=Rubin|last2=Clever|first2=H. Lawrence|journal=Chemical Reviews|volume=66|issue=4|pages=395–463}}
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| * [http://www.webqc.org/henry_gas_law.html Henry's gas law calculator]
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| * {{Cite journal|doi=10.1021/je00033a018|title=Setchenov salt-effect parameter|year=1983|last1=Clever|first1=H. Lawrence|journal=Journal of Chemical & Engineering Data|volume=28|issue=3|pages=340–343}}
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| * {{cite journal|url=http://www.nist.gov/data/PDFfiles/jpcrd606.pdf|doi=10.1063/1.1391426|title=IUPAC-NIST Solubility Data Series. 75. Nonmetals in Liquid Alkali Metals|year=2001|last1=Borgstedt|first1=Hans Ulrich|journal=Journal of Physical and Chemical Reference Data|volume=30|issue=4|pages=835–1158 }}
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| * {{Cite journal|doi=10.1063/1.443074|title=Solubilities of inert gases and methane in H2O and in D2O in the temperature range of 300 to 600 K|year=1982|last1=Crovetto|first1=Rosa|last2=Fernández‐Prini|first2=R.|last3=Japas|first3=Maria Laura|journal=J. Chem. Phys.|volume=76|issue=|page=1077}} (An example of isotopic effect on solubility)
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| {{Chemical equilibria}}
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| {{Diving medicine, physiology and physics}}
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| [[Category:Physical chemistry]]
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| [[Category:Equilibrium chemistry]]
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| [[Category:Chemical engineering]]
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| [[Category:Gas laws]]
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| [[Category:Underwater diving physics]]
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