In physical chemistry, Henry's law is a gas law that states that the amount of dissolved gas in a liquid is directly proportional at equilibrium to its partial pressure above the liquid. The proportionality factor is called Henry's law constant. It was formulated by the English chemist William Henry, who studied the topic in the early 19th century.In simple words, we can say that the partial pressure of a gas in vapour phase is directly proportional to the mole fraction of a gas in solution.
An example where Henry's law is at play is the depth-dependent dissolution of oxygen and nitrogen in the blood of underwater divers that changes during decompression, going to decompression sickness. An everyday example is 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 the solution.
In his 1803 publication about the quantity of gases absorbed by water,[1] William Henry described the results of his experiments:
Charles Coulston Gillispie states that John Dalton "supposed that the separation of gas particles one from another in the vapor phase bears the ratio of a small whole number to their interatomic distance in solution. Henry's law follows as a consequence if this ratio is a constant for each gas at a given temperature."[2]
Under high pressure, solubility of increases. On opening a container of a carbonated beverage under pressure, pressure decreases to atmospheric, so that solubility decreases and the carbon dioxide forms bubbles that are released from the liquid.
It is often noted that beer served by gravity (that is, directly from a tap in the cask) is less heavily carbonated than the same beer served via a hand-pump (or beer-engine). This is because beer is pressurised on its way to the point of service by the action of the beer engine, causing carbon dioxide to dissolve in the beer. This then comes out of solution once the beer has left the pump, causing a higher level of perceptible 'condition' in the beer.
Concentration of in the blood and tissues is so low that they feel weak and are unable to think properly, a condition called hypoxia.
In underwater diving, gas is breathed at the ambient pressure which increases with depth due to the hydrostatic pressure. Solubility of gases increases with greater depth (greater pressure) according to Henry's law, so the body tissues take on more gas over time in greater depths of water. When ascending the diver is decompressed and the solubility of the gases dissolved in the tissues decreases accordingly. If the supersaturation is too great, bubbles may form and grow, and the presence of these bubbles can cause blockages in capillaries, or distortion in the more solid tissues which can cause damage known as decompression sickness. To avoid this injury the diver must ascend slowly enough that the excess dissolved gas is carried away by the blood and released into the lung gas.
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"). This results in the Henry's law solubility constant
H\rm
H\rm
H\rm
c\rm
b
x
c\rm
p
y
c\rm
y/c\rm
cp | |
H | |
\rms |
c/p
Atmospheric chemists often define the Henry solubility as
cp | |
H | |
\rms |
=
ca | |
p |
ca
p
The SI unit for
cp | |
H | |
\rms |
ca
p
The Henry solubility can also be expressed as the dimensionless ratio between the aqueous-phase concentration
ca
cg
cc | |
H | |
\rms |
=
ca | |
cg |
For an ideal gas, the conversion is:
cc | |
H | |
\rms |
=
cp | |
RTH | |
\rms |
,
where
R
T
Sometimes, this dimensionless constant is called the water–air partitioning coefficient
KWA
L
Another Henry's law solubility constant is:
xp | |
H | |
\rms |
=
x | |
p |
Here
x
x
ca
ca ≈ x
| ||||||
|
where
\varrho | ||
|
M | ||
|
xp | |
H | |
\rms |
≈
| ||||||
|
cp | |
H | |
\rms |
The SI unit for
xp | |
H | |
\rms |
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
T
c
T
bp | |
H | |
\rms |
=
b | |
p |
.
Here
b
m
m
bp | |
H | |
\rms |
cp | |
H | |
\rms |
bp | |
H | |
\rms |
ca
b
n
i=1,\ldots,n
ca=
b\varrho | ||||||||
|
,
where
\varrho
Mi
b
bi
ca=
b\varrho | |
1+bM |
.
Henry's law is only valid for dilute solutions where
bM\ll1
\varrho ≈
\varrho | ||
|
ca ≈ b
\varrho | ||
|
and thus
bp | |
H | |
\rms |
≈
| |||||||
|
.
According to Sazonov and Shaw,[6] the dimensionless Bunsen coefficient
\alpha
cp | |
H | |
\rms |
cp | |
H | |
\rms |
=\alpha
1 | |
RTSTP |
with
TSTP
According to Sazonov and Shaw,[6] the Kuenen coefficient
S
cp | |
H | |
\rms |
cp | |
H | |
\rms |
=S
\varrho | |
RTSTP |
where
\varrho
TSTP
S
A common way to define a Henry volatility is dividing the partial pressure by the aqueous-phase concentration:
pc | |
H | |
\rmv |
=
p | |
ca |
=
1 | ||||||
|
.
The SI unit for
pc | |
H | |
\rmv |
Another Henry volatility is
px | |
H | |
\rmv |
=
p | |
x |
=
1 | ||||||
|
.
The SI unit for
px | |
H | |
\rmv |
The Henry volatility can also be expressed as the dimensionless ratio between the gas-phase concentration
cg
cc | |
H | |
\rmv |
=
cg | |
ca |
=
1 | ||||||
|
.
In chemical engineering and environmental chemistry, this dimensionless constant is often called the air–water partitioning coefficient [7] [8]
A large compilation of Henry's law constants has been published by Sander (2023). A few selected values are shown in the table below:
Gas |
=
|
=
|
=
|
=
| ||||||||||||||||||||||||||||||||||||
---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|
\right) |
\right) | \left(atm\right) | (dimensionless) | |||||||||||||||||||||||||||||||||||||
770 | 1.3 | 4.3 | 3.2 | |||||||||||||||||||||||||||||||||||||
1300 | 7.8 | 7.1 | 1.9 | |||||||||||||||||||||||||||||||||||||
29 | 3.4 | 1.6 | 8.3 | |||||||||||||||||||||||||||||||||||||
1600 | 6.1 | 9.1 | 1.5 | |||||||||||||||||||||||||||||||||||||
2700 | 3.7 | 1.5 | 9.1 | |||||||||||||||||||||||||||||||||||||
2200 | 4.5 | 1.2 | 1.1 | |||||||||||||||||||||||||||||||||||||
710 | 1.4 | 4.0 | 3.4 | |||||||||||||||||||||||||||||||||||||
1100 | 9.5 | 5.8 | 2.3 |
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:
dlnH | |
d(1/T) |
=
-\DeltasolH | |
R |
,
where
\DeltasolH
H
\DeltasolH
H
Hs
Hv
Integrating the above equation and creating an expression based on
H\circ
T\circ
H(T)=
| |||||
H | \left( |
1 | |
T |
-
1 | |
T\circ |
\right)\right].
The van 't Hoff equation in this form is only valid for a limited temperature range in which
\DeltasolH
The following table lists some temperature dependencies:
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.[10]
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, formaldehyde is almost completely hydrated:
The total concentration of dissolved formaldehyde is
Taking this equilibrium into account, an effective Henry's law constant
H\rm
For acids and bases, the effective Henry's law constant is not a useful quantity because it depends on the pH of the solution. In order to obtain a pH-independent constant, the product of the intrinsic Henry's law constant and the acidity constant is often used for strong acids like hydrochloric acid (HCl):
Although
H'
Values of Henry's law constants for aqueous solutions depend on the composition of the solution, i.e., on its ionic strength and on dissolved organics. In general, the solubility of a gas decreases with increasing salinity ("salting out"). However, a "salting in" effect has also been observed, for example for the effective Henry's law constant of glyoxal. The effect can be described with the Sechenov equation, named after the Russian physiologist Ivan Sechenov (sometimes the German transliteration "Setschenow" of the Cyrillic name Се́ченов is used). There are many alternative ways to define the Sechenov equation, depending on how the aqueous-phase composition is described (based on concentration, molality, or molar fraction) and which variant of the Henry's law constant is used. Describing the solution in terms of molality is preferred because molality is invariant to temperature and to the addition of dry salt to the solution. Thus, the Sechenov equation can be written as
log\left( |
| ||||||
|
\right)=k\rmb(salt),
where
bp | |
H | |
\rms,0 |
bp | |
H | |
\rms |
k\rm
b(salt)
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. In these cases, it is necessary to state the law in terms of chemical potentials. For a solute in an ideal dilute solution, the chemical potential depends only on the concentration. For non-ideal solutions, the activity coefficients of the components must be taken into account:
\mu=
\circ | |
\mu | |
c |
+RTln
\gammacc | |
c\circ |
where
\gammac=
H\rm | |
p* |
For non-ideal solutions, the infinite dilution activity coefficient γc 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:
\suminid\mui=0.
By measuring the change in vapor pressure (and hence chemical potential) of the solvent, the chemical potential of the solute can be deduced.
The standard state for a dilute solution is also defined in terms of infinite-dilution behavior. Although the standard concentration c° 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).
All the relations above can also be expressed in terms of molalities b rather than concentrations, e.g.:
\mu=
\circ | |
\mu | |
b |
+RTln
\gammabb | |
b\circ |
,
where
\gammab=
| |||||||
p* |
The standard chemical potential μm°, the activity coefficient γm and the Henry's law constant Hvpb all have different numerical values when molalities are used in place of concentrations.
Henry's law solubility constant
xp | |
H | |
\rms,2,M |
xp | |
H | |
\rms,2,1 |
xp | |
H | |
\rms,2,3 |
ln
xp | |
H | |
\rms,2,M |
=x1ln
xp | |
H | |
\rms,2,1 |
+x3ln
xp | |
H | |
\rms,2,3 |
+a13x1x3
Where
x1
x3
A similar relationship can be found for the volatility constant
px | |
H | |
\rmv,2,M |
px | |
H | |
\rmv |
xp | |
=1/H | |
\rms |
ln
xp | |
H | |
\rms |
=-ln
xp | |
(1/H | |
\rms |
)=-ln
px | |
H | |
\rmv |
ln
px | |
H | |
\rmv,2,M |
=x1ln
px | |
H | |
\rmv,2,1 |
+x3ln
px | |
H | |
\rmv,2,3 |
-a13x1x3
For a water-ethanol mixture, the interaction parameter a13 has values around
0.1\pm0.05
In geochemistry, a version of Henry's law applies to the solubility of a noble gas in contact with silicate melt. One equation used is
Cmelt | |
Cgas |
=
E | |
\exp\left[-\beta\left(\mu | |
melt |
-
E | |
\mu | |
gas\right)\right], |
where
C is the number concentrations of the solute gas in the melt and gas phases,
β = 1/kBT, an inverse temperature parameter (kB is the Boltzmann constant),
μE is the excess chemical potentials of the solute gas in the two phases.
Henry's law is a limiting law that only applies for "sufficiently dilute" solutions, while Raoult's law is generally valid when the liquid phase is almost pure or for mixtures of similar substances.[13] The range of concentrations in which Henry's law 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.
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
p=
px | |
H | |
\rmv |
x.
This can be compared with Raoult's law:
p=p*x,
where p* is the vapor pressure of the pure component.
At first sight, Raoult's law appears to be a special case of Henry's law, where Hvpx = 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.
The general case is that both laws are 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:
Raoult's law:
\limx\left(
p | |
x |
\right)=p*.
Henry's law:
\limx\left(
p | |
x |
\right)=
px | |
H | |
\rmv |
.
Raoult's law can also be related to non-gas solutes.