In physical chemistry, Henry's law is a gas law that states that the amount of dissolved gas in a liquid is proportional 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 his publication about the quantity of gases absorbed by water, 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.
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 . 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-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. 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
The SI unit for is mol/(m3·Pa); however, often the unit M/atm is used, since is usually expressed in M (1M = 1 mol/dm3) and in atm (1atm = 101325Pa).
The Henry solubility can also be expressed as the dimensionless ratio between the aqueous-phase concentration of a species and its gas-phase concentration :
For an ideal gas, the conversion is
where is the gas constant, and is the temperature.
Sometimes, this dimensionless constant is called the water-air partitioning coefficient . It is closely related to the various, slightly different definitions of the Ostwald coefficient , as discussed by Battino (1984).
Another Henry's law solubility constant is
Here is the molar mixing ratio in the aqueous phase. For a dilute aqueous solution the conversion between and is:
where is the density of water and is the molar mass of water. Thus
The SI unit for is Pa-1, although atm-1 is still frequently used.
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
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 V2* of pure solvent at the temperature of measurement and partial pressure of 1 bar." If the gas is ideal, the pressure cancels out, and the conversion to is simply
with = 273.15K. Note, that according to this definition, the conversion factor is not temperature-dependent. Independent of the temperature that the Bunsen coefficient refers to, 273.15K 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." 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 m3/kg. The Kuenen coefficient, which is named after Johannes Kuenen, has been used mainly in the older literature, and IUPAC considers it to be obsolete.
A common way to define a Henry volatility is dividing the partial pressure by the aqueous-phase concentration:
The SI unit for is Pa·m3/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 :
A large compilation of Henry's law constants has been published by Sander (2015). A few selected values are shown in the table below:
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:
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.
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 can be defined as
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 is usually also called a Henry's law constant, it is a different quantity and it has different units than .
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 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
where is the Henry's law constant in pure water, is the Henry's law constant in the salt solution, is the molality-based Sechenov constant, and is the molality of the 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:
where for a volatile solute; c° = 1 mol/L.
For non-ideal solutions, the activity coefficient ?c 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:
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.:
where for a volatile solute; b° = 1 mol/kg.
The standard chemical potential ?m°, the activity coefficient ?m and the Henry's law constant KH,b all have different numerical values when molalities are used in place of concentrations.
Henry law constant H2, M for a gas 2 in a mixture of solvents 1 and 3 is related to the constants for individual solvents H21 and H23:
where a13 is the interaction parameter of the solvents from Wohl expansion of the excess chemical potential of the ternary mixtures.
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. 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
This can be compared with Raoult's law:
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 KH = 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 can also be related to non-gas solutes.