Henry's law- Solutions, CBSE, Class 12, Chemistry PDF Download

 Page 1


Henry's law
In chemistry, Henry's law is a gas law that states that the amount of dissolved gas is proportional to its partial pressure in the gas
phase. The proportionality factor is called the Henry's law constant. It was formulated by the English chemist William Henry, who
studied the topic in the early 19th century . In his publication about the quantity of gases absorbed by water,
[1]
 he described the results
of his experiments:
..."water takes up, of gas condensed by one, two, or more additional atmospheres, a quantity
which, ordinarily compressed, would be equal to twice, thrice, &c. the volume absorbed
under the common pressure of the atmosphere."
An example where Henry's law is at play is in the depth-dependent dissolution of oxygen and nitrogen in the blood of underwater
divers that changes during decompression , leading to decompression sickness. An everyday example is given by one's experience
with carbonated soft drinks, which contain dissolved carbon dioxide. Before opening, the gas above the drink in its container is
almost pure carbon dioxide , at a pressure higher than atmospheric pressure . After the bottle is opened, this gas escapes, moving the
partial pressure of carbon dioxide above the liquid to be much lower, resulting in degassing as the dissolved carbon dioxide comes
out of solution.
Fundamental types and variants of Henry's law constants
Henry's law solubility constants 
Henry's law volatility constants 
Values of Henry's law constants
Temperature dependence
Effective Henry's law constants H
eff
Dependence on ionic strength (Sechenov equation)
Non-ideal solutions
Solvent mixtures
Miscellaneous
In geochemistry
Comparison to Raoult's law
See also
References
External links
There are many ways to define the proportionality constant of Henry's law, which can be subdivided into two fundamental types: One
possibility is to put the aqueous phase into the numerator and the gaseous phase into the denominator ("aq/gas").
[2]
 This results in the
Henry's law solubility constant . Its value increases with increased solubility . Alternatively , numerator and denominator can be
switched ("gas/aq"), which results in the Henry's law volatility constant . The value of  decreases with increased solubility .
There are several variants of both fundamental types. This results from the multiplicity of quantities that can be chosen to describe
the composition of the two phases. Typical choices for the aqueous phase are molar concentration ( ), molality ( ), and molar
mixing ratio ( ). For the gas phase, molar concentration ( ) and partial pressure ( ) are often used. It is not possible to use the gas-
Contents
Fundamental types and variants of Henry's law constants
Page 2


Henry's law
In chemistry, Henry's law is a gas law that states that the amount of dissolved gas is proportional to its partial pressure in the gas
phase. The proportionality factor is called the Henry's law constant. It was formulated by the English chemist William Henry, who
studied the topic in the early 19th century . In his publication about the quantity of gases absorbed by water,
[1]
 he described the results
of his experiments:
..."water takes up, of gas condensed by one, two, or more additional atmospheres, a quantity
which, ordinarily compressed, would be equal to twice, thrice, &c. the volume absorbed
under the common pressure of the atmosphere."
An example where Henry's law is at play is in the depth-dependent dissolution of oxygen and nitrogen in the blood of underwater
divers that changes during decompression , leading to decompression sickness. An everyday example is given by one's experience
with carbonated soft drinks, which contain dissolved carbon dioxide. Before opening, the gas above the drink in its container is
almost pure carbon dioxide , at a pressure higher than atmospheric pressure . After the bottle is opened, this gas escapes, moving the
partial pressure of carbon dioxide above the liquid to be much lower, resulting in degassing as the dissolved carbon dioxide comes
out of solution.
Fundamental types and variants of Henry's law constants
Henry's law solubility constants 
Henry's law volatility constants 
Values of Henry's law constants
Temperature dependence
Effective Henry's law constants H
eff
Dependence on ionic strength (Sechenov equation)
Non-ideal solutions
Solvent mixtures
Miscellaneous
In geochemistry
Comparison to Raoult's law
See also
References
External links
There are many ways to define the proportionality constant of Henry's law, which can be subdivided into two fundamental types: One
possibility is to put the aqueous phase into the numerator and the gaseous phase into the denominator ("aq/gas").
[2]
 This results in the
Henry's law solubility constant . Its value increases with increased solubility . Alternatively , numerator and denominator can be
switched ("gas/aq"), which results in the Henry's law volatility constant . The value of  decreases with increased solubility .
There are several variants of both fundamental types. This results from the multiplicity of quantities that can be chosen to describe
the composition of the two phases. Typical choices for the aqueous phase are molar concentration ( ), molality ( ), and molar
mixing ratio ( ). For the gas phase, molar concentration ( ) and partial pressure ( ) are often used. It is not possible to use the gas-
Contents
Fundamental types and variants of Henry's law constants
phase mixing ratio ( ) because at a given gas-phase mixing ratio, the aqueous-phase concentration  depends on the total pressure
and thus the ratio  is not a constant.
[3]
 To specify the exact variant of the Henry's law constant, two superscripts are used. They
refer to the numerator and the denominator of the definition. For example,  refers to the Henry solubility defined as .
Atmospheric chemists often define the Henry solubility as
.
[2]
Here  is the concentration of a species in the aqueous phase, and  is the partial pressure of that species in the gas phase under
equilibrium conditions.
The SI unit for  is mol/(m
3
 Pa); however , often the unit M/atm is used, since  is usually expressed in M (1 M = 1 mol/dm
3
) and
 in atm (1 atm = 101325 Pa).
The Henry solubility can also be expressed as the dimensionless ratio between the aqueous-phase concentration  of a species and
its gas-phase concentration :
.
[2]
For an ideal gas, the conversion is:
 ,
[2]
where  is the gas constant and  is the temperature.
Sometimes, this dimensionless constant is called the "water-air partitioning coefficient" .
[4]
 It is closely related to the various,
slightly different definitions of the "Ostwald coef ficient" , as discussed by Battino (1984).
[5]
Another Henry's law solubility constant is
 .
[2]
Here  is the molar mixing ratio in the aqueous phase. For a dilute aqueous solution the conversion between  and  is:
 ,
[2]
where  is the density of water and  is the molar mass of water . Thus
 .
[2]
The SI unit for  is Pa
-1
, although atm
-1
 is still frequently used.
[2]
Henry's law solubility constants 
Henry solubility defined via concentration ( )
The dimensionless Henry solubility 
Henry solubility defined via aqueous-phase mixing ratio ( )
Henry solubility defined via molality ( )
Page 3


Henry's law
In chemistry, Henry's law is a gas law that states that the amount of dissolved gas is proportional to its partial pressure in the gas
phase. The proportionality factor is called the Henry's law constant. It was formulated by the English chemist William Henry, who
studied the topic in the early 19th century . In his publication about the quantity of gases absorbed by water,
[1]
 he described the results
of his experiments:
..."water takes up, of gas condensed by one, two, or more additional atmospheres, a quantity
which, ordinarily compressed, would be equal to twice, thrice, &c. the volume absorbed
under the common pressure of the atmosphere."
An example where Henry's law is at play is in the depth-dependent dissolution of oxygen and nitrogen in the blood of underwater
divers that changes during decompression , leading to decompression sickness. An everyday example is given by one's experience
with carbonated soft drinks, which contain dissolved carbon dioxide. Before opening, the gas above the drink in its container is
almost pure carbon dioxide , at a pressure higher than atmospheric pressure . After the bottle is opened, this gas escapes, moving the
partial pressure of carbon dioxide above the liquid to be much lower, resulting in degassing as the dissolved carbon dioxide comes
out of solution.
Fundamental types and variants of Henry's law constants
Henry's law solubility constants 
Henry's law volatility constants 
Values of Henry's law constants
Temperature dependence
Effective Henry's law constants H
eff
Dependence on ionic strength (Sechenov equation)
Non-ideal solutions
Solvent mixtures
Miscellaneous
In geochemistry
Comparison to Raoult's law
See also
References
External links
There are many ways to define the proportionality constant of Henry's law, which can be subdivided into two fundamental types: One
possibility is to put the aqueous phase into the numerator and the gaseous phase into the denominator ("aq/gas").
[2]
 This results in the
Henry's law solubility constant . Its value increases with increased solubility . Alternatively , numerator and denominator can be
switched ("gas/aq"), which results in the Henry's law volatility constant . The value of  decreases with increased solubility .
There are several variants of both fundamental types. This results from the multiplicity of quantities that can be chosen to describe
the composition of the two phases. Typical choices for the aqueous phase are molar concentration ( ), molality ( ), and molar
mixing ratio ( ). For the gas phase, molar concentration ( ) and partial pressure ( ) are often used. It is not possible to use the gas-
Contents
Fundamental types and variants of Henry's law constants
phase mixing ratio ( ) because at a given gas-phase mixing ratio, the aqueous-phase concentration  depends on the total pressure
and thus the ratio  is not a constant.
[3]
 To specify the exact variant of the Henry's law constant, two superscripts are used. They
refer to the numerator and the denominator of the definition. For example,  refers to the Henry solubility defined as .
Atmospheric chemists often define the Henry solubility as
.
[2]
Here  is the concentration of a species in the aqueous phase, and  is the partial pressure of that species in the gas phase under
equilibrium conditions.
The SI unit for  is mol/(m
3
 Pa); however , often the unit M/atm is used, since  is usually expressed in M (1 M = 1 mol/dm
3
) and
 in atm (1 atm = 101325 Pa).
The Henry solubility can also be expressed as the dimensionless ratio between the aqueous-phase concentration  of a species and
its gas-phase concentration :
.
[2]
For an ideal gas, the conversion is:
 ,
[2]
where  is the gas constant and  is the temperature.
Sometimes, this dimensionless constant is called the "water-air partitioning coefficient" .
[4]
 It is closely related to the various,
slightly different definitions of the "Ostwald coef ficient" , as discussed by Battino (1984).
[5]
Another Henry's law solubility constant is
 .
[2]
Here  is the molar mixing ratio in the aqueous phase. For a dilute aqueous solution the conversion between  and  is:
 ,
[2]
where  is the density of water and  is the molar mass of water . Thus
 .
[2]
The SI unit for  is Pa
-1
, although atm
-1
 is still frequently used.
[2]
Henry's law solubility constants 
Henry solubility defined via concentration ( )
The dimensionless Henry solubility 
Henry solubility defined via aqueous-phase mixing ratio ( )
Henry solubility defined via molality ( )
It can be advantageous to describe the aqueous phase in terms of molality instead of concentration. The molality of a solution does
not change with , since it refers to the mass of the solvent. In contrast, the concentration  does change with , since the density of
a solution and thus its volume are temperature-dependent. Defining the aqueous-phase composition via molality has the advantage
that any temperature dependence of the Henry's law constant is a true solubility phenomenon and not introduced indirectly via a
density change of the solution. Using molality , the Henry solubility can be defined as
Here  is used as the symbol for molality (instead of ) to avoid confusion with the symbol  for mass. The SI unit for  is
mol/(kg Pa). There is no simple way to calculate  from , since the conversion between concentration  and molality 
involves all solutes of a solution. For a solution with a total of  solutes with indices , the conversion is:
where  is the density of the solution, and  are the molar masses. Here  is identical to one of the  in the denominator. If there is
only one solute, the equation simplifies to
Henry's law is only valid for dilute solutions where  and . In this case the conversion reduces further to
and thus
According to Sazonov and Shaw, the dimensionless Bunsen coefficient  is defined as "the volume of saturating gas, V1, reduced to
T° = 273.15 K, p° = 1 bar, which is absorbed by unit volume V
2
* of pure solvent at the temperature of measurement and partial
pressure of 1 bar ."
[6]
 If the gas is ideal, the pressure cancels out, and the conversion to  is simply
 ,
with  = 273.15 K. Note, that according to this definition, the conversion factor is not temperature-dependent. Independent of the
temperature that the Bunsen coefficient refers to, 273.15 K is always used for the conversion. The Bunsen coefficient, which is
named after Robert Bunsen , has been used mainly in the older literature.
According to Sazonov and Shaw, the Kuenen coefficient  is defined as "the volume of saturating gas V(g), reduced to T° = 273.15
K, p° = bar, which is dissolved by unit mass of pure solvent at the temperature of measurement and partial pressure 1 bar."
[6]
 If the
gas is ideal, the relation to  is
 ,
where  is the density of the solvent, and  = 273.15 K. The SI unit for  is m
3
/kg.
[6]
 The Kuenen coefficient, which is named
after Johannes Kuenen , has been used mainly in the older literature, and IUP AC considers it to be obsolete.
[7]
The Bunsen coefficient 
The Kuenen coefficient 
Page 4


Henry's law
In chemistry, Henry's law is a gas law that states that the amount of dissolved gas is proportional to its partial pressure in the gas
phase. The proportionality factor is called the Henry's law constant. It was formulated by the English chemist William Henry, who
studied the topic in the early 19th century . In his publication about the quantity of gases absorbed by water,
[1]
 he described the results
of his experiments:
..."water takes up, of gas condensed by one, two, or more additional atmospheres, a quantity
which, ordinarily compressed, would be equal to twice, thrice, &c. the volume absorbed
under the common pressure of the atmosphere."
An example where Henry's law is at play is in the depth-dependent dissolution of oxygen and nitrogen in the blood of underwater
divers that changes during decompression , leading to decompression sickness. An everyday example is given by one's experience
with carbonated soft drinks, which contain dissolved carbon dioxide. Before opening, the gas above the drink in its container is
almost pure carbon dioxide , at a pressure higher than atmospheric pressure . After the bottle is opened, this gas escapes, moving the
partial pressure of carbon dioxide above the liquid to be much lower, resulting in degassing as the dissolved carbon dioxide comes
out of solution.
Fundamental types and variants of Henry's law constants
Henry's law solubility constants 
Henry's law volatility constants 
Values of Henry's law constants
Temperature dependence
Effective Henry's law constants H
eff
Dependence on ionic strength (Sechenov equation)
Non-ideal solutions
Solvent mixtures
Miscellaneous
In geochemistry
Comparison to Raoult's law
See also
References
External links
There are many ways to define the proportionality constant of Henry's law, which can be subdivided into two fundamental types: One
possibility is to put the aqueous phase into the numerator and the gaseous phase into the denominator ("aq/gas").
[2]
 This results in the
Henry's law solubility constant . Its value increases with increased solubility . Alternatively , numerator and denominator can be
switched ("gas/aq"), which results in the Henry's law volatility constant . The value of  decreases with increased solubility .
There are several variants of both fundamental types. This results from the multiplicity of quantities that can be chosen to describe
the composition of the two phases. Typical choices for the aqueous phase are molar concentration ( ), molality ( ), and molar
mixing ratio ( ). For the gas phase, molar concentration ( ) and partial pressure ( ) are often used. It is not possible to use the gas-
Contents
Fundamental types and variants of Henry's law constants
phase mixing ratio ( ) because at a given gas-phase mixing ratio, the aqueous-phase concentration  depends on the total pressure
and thus the ratio  is not a constant.
[3]
 To specify the exact variant of the Henry's law constant, two superscripts are used. They
refer to the numerator and the denominator of the definition. For example,  refers to the Henry solubility defined as .
Atmospheric chemists often define the Henry solubility as
.
[2]
Here  is the concentration of a species in the aqueous phase, and  is the partial pressure of that species in the gas phase under
equilibrium conditions.
The SI unit for  is mol/(m
3
 Pa); however , often the unit M/atm is used, since  is usually expressed in M (1 M = 1 mol/dm
3
) and
 in atm (1 atm = 101325 Pa).
The Henry solubility can also be expressed as the dimensionless ratio between the aqueous-phase concentration  of a species and
its gas-phase concentration :
.
[2]
For an ideal gas, the conversion is:
 ,
[2]
where  is the gas constant and  is the temperature.
Sometimes, this dimensionless constant is called the "water-air partitioning coefficient" .
[4]
 It is closely related to the various,
slightly different definitions of the "Ostwald coef ficient" , as discussed by Battino (1984).
[5]
Another Henry's law solubility constant is
 .
[2]
Here  is the molar mixing ratio in the aqueous phase. For a dilute aqueous solution the conversion between  and  is:
 ,
[2]
where  is the density of water and  is the molar mass of water . Thus
 .
[2]
The SI unit for  is Pa
-1
, although atm
-1
 is still frequently used.
[2]
Henry's law solubility constants 
Henry solubility defined via concentration ( )
The dimensionless Henry solubility 
Henry solubility defined via aqueous-phase mixing ratio ( )
Henry solubility defined via molality ( )
It can be advantageous to describe the aqueous phase in terms of molality instead of concentration. The molality of a solution does
not change with , since it refers to the mass of the solvent. In contrast, the concentration  does change with , since the density of
a solution and thus its volume are temperature-dependent. Defining the aqueous-phase composition via molality has the advantage
that any temperature dependence of the Henry's law constant is a true solubility phenomenon and not introduced indirectly via a
density change of the solution. Using molality , the Henry solubility can be defined as
Here  is used as the symbol for molality (instead of ) to avoid confusion with the symbol  for mass. The SI unit for  is
mol/(kg Pa). There is no simple way to calculate  from , since the conversion between concentration  and molality 
involves all solutes of a solution. For a solution with a total of  solutes with indices , the conversion is:
where  is the density of the solution, and  are the molar masses. Here  is identical to one of the  in the denominator. If there is
only one solute, the equation simplifies to
Henry's law is only valid for dilute solutions where  and . In this case the conversion reduces further to
and thus
According to Sazonov and Shaw, the dimensionless Bunsen coefficient  is defined as "the volume of saturating gas, V1, reduced to
T° = 273.15 K, p° = 1 bar, which is absorbed by unit volume V
2
* of pure solvent at the temperature of measurement and partial
pressure of 1 bar ."
[6]
 If the gas is ideal, the pressure cancels out, and the conversion to  is simply
 ,
with  = 273.15 K. Note, that according to this definition, the conversion factor is not temperature-dependent. Independent of the
temperature that the Bunsen coefficient refers to, 273.15 K is always used for the conversion. The Bunsen coefficient, which is
named after Robert Bunsen , has been used mainly in the older literature.
According to Sazonov and Shaw, the Kuenen coefficient  is defined as "the volume of saturating gas V(g), reduced to T° = 273.15
K, p° = bar, which is dissolved by unit mass of pure solvent at the temperature of measurement and partial pressure 1 bar."
[6]
 If the
gas is ideal, the relation to  is
 ,
where  is the density of the solvent, and  = 273.15 K. The SI unit for  is m
3
/kg.
[6]
 The Kuenen coefficient, which is named
after Johannes Kuenen , has been used mainly in the older literature, and IUP AC considers it to be obsolete.
[7]
The Bunsen coefficient 
The Kuenen coefficient 
A common way to define a Henry volatility is dividing the partial pressure by the aqueous-phase concentration:
The SI unit for  is Pa m
3
/mol.
Another Henry volatility is
The SI unit for  is Pa. However , atm is still frequently used.
The Henry volatility can also be expressed as the dimensionless ratio between the gas-phase concentration  of a species and its
aqueous-phase concentration :
In chemical engineering and environmental chemistry, this dimensionless constant is often called the air–water partitioning
coefficient .
A large compilation of Henry's law constants has been published by Sander (2015).
[2]
 A few selected values are shown in the table
below:
Henry's law constants (gases in water at 298.15 K)
equation:
unit: (dimensionless)
O
2
770 1.3 ×10
-3
4.3 ×10
4
3.2 ×10
-2
H
2
1300 7.8 ×10
-4
7.1 ×10
4
1.9 ×10
-2
CO
2
29 3.4 ×10
-2
1.6 ×10
3
8.3 ×10
-1
N
2
1600 6.1 ×10
-4
9.1 ×10
4
1.5 ×10
-2
He 2700 3.7 ×10
-4
1.5 ×10
5
9.1 ×10
-3
Ne 2200 4.5 ×10
-4
1.2 ×10
5
1.1 ×10
-2
Ar 710 1.4 ×10
-3
4.0 ×10
4
3.4 ×10
-2
CO 1100 9.5 ×10
-4
5.8 ×10
4
2.3 ×10
-2
Henry's law volatility constants 
The Henry volatility defined via concentration ( )
The Henry volatility defined via aqueous-phase mixing ratio ( )
The dimensionless Henry volatility 
Values of Henry's law constants
Temperature dependence
Page 5


Henry's law
In chemistry, Henry's law is a gas law that states that the amount of dissolved gas is proportional to its partial pressure in the gas
phase. The proportionality factor is called the Henry's law constant. It was formulated by the English chemist William Henry, who
studied the topic in the early 19th century . In his publication about the quantity of gases absorbed by water,
[1]
 he described the results
of his experiments:
..."water takes up, of gas condensed by one, two, or more additional atmospheres, a quantity
which, ordinarily compressed, would be equal to twice, thrice, &c. the volume absorbed
under the common pressure of the atmosphere."
An example where Henry's law is at play is in the depth-dependent dissolution of oxygen and nitrogen in the blood of underwater
divers that changes during decompression , leading to decompression sickness. An everyday example is given by one's experience
with carbonated soft drinks, which contain dissolved carbon dioxide. Before opening, the gas above the drink in its container is
almost pure carbon dioxide , at a pressure higher than atmospheric pressure . After the bottle is opened, this gas escapes, moving the
partial pressure of carbon dioxide above the liquid to be much lower, resulting in degassing as the dissolved carbon dioxide comes
out of solution.
Fundamental types and variants of Henry's law constants
Henry's law solubility constants 
Henry's law volatility constants 
Values of Henry's law constants
Temperature dependence
Effective Henry's law constants H
eff
Dependence on ionic strength (Sechenov equation)
Non-ideal solutions
Solvent mixtures
Miscellaneous
In geochemistry
Comparison to Raoult's law
See also
References
External links
There are many ways to define the proportionality constant of Henry's law, which can be subdivided into two fundamental types: One
possibility is to put the aqueous phase into the numerator and the gaseous phase into the denominator ("aq/gas").
[2]
 This results in the
Henry's law solubility constant . Its value increases with increased solubility . Alternatively , numerator and denominator can be
switched ("gas/aq"), which results in the Henry's law volatility constant . The value of  decreases with increased solubility .
There are several variants of both fundamental types. This results from the multiplicity of quantities that can be chosen to describe
the composition of the two phases. Typical choices for the aqueous phase are molar concentration ( ), molality ( ), and molar
mixing ratio ( ). For the gas phase, molar concentration ( ) and partial pressure ( ) are often used. It is not possible to use the gas-
Contents
Fundamental types and variants of Henry's law constants
phase mixing ratio ( ) because at a given gas-phase mixing ratio, the aqueous-phase concentration  depends on the total pressure
and thus the ratio  is not a constant.
[3]
 To specify the exact variant of the Henry's law constant, two superscripts are used. They
refer to the numerator and the denominator of the definition. For example,  refers to the Henry solubility defined as .
Atmospheric chemists often define the Henry solubility as
.
[2]
Here  is the concentration of a species in the aqueous phase, and  is the partial pressure of that species in the gas phase under
equilibrium conditions.
The SI unit for  is mol/(m
3
 Pa); however , often the unit M/atm is used, since  is usually expressed in M (1 M = 1 mol/dm
3
) and
 in atm (1 atm = 101325 Pa).
The Henry solubility can also be expressed as the dimensionless ratio between the aqueous-phase concentration  of a species and
its gas-phase concentration :
.
[2]
For an ideal gas, the conversion is:
 ,
[2]
where  is the gas constant and  is the temperature.
Sometimes, this dimensionless constant is called the "water-air partitioning coefficient" .
[4]
 It is closely related to the various,
slightly different definitions of the "Ostwald coef ficient" , as discussed by Battino (1984).
[5]
Another Henry's law solubility constant is
 .
[2]
Here  is the molar mixing ratio in the aqueous phase. For a dilute aqueous solution the conversion between  and  is:
 ,
[2]
where  is the density of water and  is the molar mass of water . Thus
 .
[2]
The SI unit for  is Pa
-1
, although atm
-1
 is still frequently used.
[2]
Henry's law solubility constants 
Henry solubility defined via concentration ( )
The dimensionless Henry solubility 
Henry solubility defined via aqueous-phase mixing ratio ( )
Henry solubility defined via molality ( )
It can be advantageous to describe the aqueous phase in terms of molality instead of concentration. The molality of a solution does
not change with , since it refers to the mass of the solvent. In contrast, the concentration  does change with , since the density of
a solution and thus its volume are temperature-dependent. Defining the aqueous-phase composition via molality has the advantage
that any temperature dependence of the Henry's law constant is a true solubility phenomenon and not introduced indirectly via a
density change of the solution. Using molality , the Henry solubility can be defined as
Here  is used as the symbol for molality (instead of ) to avoid confusion with the symbol  for mass. The SI unit for  is
mol/(kg Pa). There is no simple way to calculate  from , since the conversion between concentration  and molality 
involves all solutes of a solution. For a solution with a total of  solutes with indices , the conversion is:
where  is the density of the solution, and  are the molar masses. Here  is identical to one of the  in the denominator. If there is
only one solute, the equation simplifies to
Henry's law is only valid for dilute solutions where  and . In this case the conversion reduces further to
and thus
According to Sazonov and Shaw, the dimensionless Bunsen coefficient  is defined as "the volume of saturating gas, V1, reduced to
T° = 273.15 K, p° = 1 bar, which is absorbed by unit volume V
2
* of pure solvent at the temperature of measurement and partial
pressure of 1 bar ."
[6]
 If the gas is ideal, the pressure cancels out, and the conversion to  is simply
 ,
with  = 273.15 K. Note, that according to this definition, the conversion factor is not temperature-dependent. Independent of the
temperature that the Bunsen coefficient refers to, 273.15 K is always used for the conversion. The Bunsen coefficient, which is
named after Robert Bunsen , has been used mainly in the older literature.
According to Sazonov and Shaw, the Kuenen coefficient  is defined as "the volume of saturating gas V(g), reduced to T° = 273.15
K, p° = bar, which is dissolved by unit mass of pure solvent at the temperature of measurement and partial pressure 1 bar."
[6]
 If the
gas is ideal, the relation to  is
 ,
where  is the density of the solvent, and  = 273.15 K. The SI unit for  is m
3
/kg.
[6]
 The Kuenen coefficient, which is named
after Johannes Kuenen , has been used mainly in the older literature, and IUP AC considers it to be obsolete.
[7]
The Bunsen coefficient 
The Kuenen coefficient 
A common way to define a Henry volatility is dividing the partial pressure by the aqueous-phase concentration:
The SI unit for  is Pa m
3
/mol.
Another Henry volatility is
The SI unit for  is Pa. However , atm is still frequently used.
The Henry volatility can also be expressed as the dimensionless ratio between the gas-phase concentration  of a species and its
aqueous-phase concentration :
In chemical engineering and environmental chemistry, this dimensionless constant is often called the air–water partitioning
coefficient .
A large compilation of Henry's law constants has been published by Sander (2015).
[2]
 A few selected values are shown in the table
below:
Henry's law constants (gases in water at 298.15 K)
equation:
unit: (dimensionless)
O
2
770 1.3 ×10
-3
4.3 ×10
4
3.2 ×10
-2
H
2
1300 7.8 ×10
-4
7.1 ×10
4
1.9 ×10
-2
CO
2
29 3.4 ×10
-2
1.6 ×10
3
8.3 ×10
-1
N
2
1600 6.1 ×10
-4
9.1 ×10
4
1.5 ×10
-2
He 2700 3.7 ×10
-4
1.5 ×10
5
9.1 ×10
-3
Ne 2200 4.5 ×10
-4
1.2 ×10
5
1.1 ×10
-2
Ar 710 1.4 ×10
-3
4.0 ×10
4
3.4 ×10
-2
CO 1100 9.5 ×10
-4
5.8 ×10
4
2.3 ×10
-2
Henry's law volatility constants 
The Henry volatility defined via concentration ( )
The Henry volatility defined via aqueous-phase mixing ratio ( )
The dimensionless Henry volatility 
Values of Henry's law constants
Temperature dependence
When the temperature of a system changes, the Henry constant also changes. The temperature dependence of equilibrium constants
can generally be described with the van 't Hoff equation, which also applies to Henry's law constants:
where  is the enthalpy of dissolution. Note that the letter  in the symbol  refers to enthalpy and is not related to the
letter  for Henry's law constants. Integrating the above equation and creating an expression based on  at the reference
temperature  = 298.15 K yields:
The van 't Hoff equation in this form is only valid for a limited temperature range in which  does not change much with
temperature.
The following table lists some temperature dependencies:
Values of  (in K)
O
2
H
2
CO
2
N
2
He Ne Ar CO
 1700  500  2400  1300  230  490  1300  1300 
Solubility of permanent gases usually decreases with increasing temperature at around room temperature. However , for aqueous
solutions, the Henry's law solubility constant for many species 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 at about 30 °C for helium, 92 to 93 °C for argon, nitrogen and oxygen,
and 114 °C for xenon.
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The Henry's law constants mentioned so far do not consider any chemical equilibria in the aqueous phase. This type is called the
"intrinsic" (or "physical") Henry's law constant. For example, the intrinsic Henry's law solubility constant of formaldehyde can be
defined as
In aqueous solution, methanal is almost completely hydrated:
The total concentration of dissolved methanal is
Taking this equilibrium into account, an ef fective Henry's law constant  can be defined as
Effective Henry's law constants H
eff