Henry's law
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.
It was formulated by the English chemist William Henry, who studied the topic in the early 19th century.
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: … 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.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.
Concentration of O2 in the blood and tissues is so low that they feel weak and are unable to think properly, a condition called hypoxia.
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.
Alternatively, numerator and denominator can be switched ("gas/aq"), which results in the Henry's law volatility constant
is the partial pressure of that species in the gas phase under equilibrium conditions.
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.
In this case the conversion reduces further to and thus According to Sazonov and Shaw,[7] the dimensionless Bunsen coefficient
[3] A common way to define a Henry volatility is dividing the partial pressure by the aqueous-phase concentration: The SI unit for
The Henry volatility can also be expressed as the dimensionless ratio between the gas-phase concentration
: In chemical engineering and environmental chemistry, this dimensionless constant is often called the air–water partitioning coefficient
[8][9] A large compilation of Henry's law constants has been published by Sander (2023).
= 298.15 K yields: The van 't Hoff equation in this form is only valid for a limited temperature range in which
However, for aqueous solutions, the Henry's law solubility constant for many species goes through a minimum.
[12] The Henry's law constants mentioned so far do not consider any chemical equilibria in the aqueous phase.
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 general, the solubility of a gas decreases with increasing salinity ("salting out").
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.
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.
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.
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.
[14] In geochemistry, a version of Henry's law applies to the solubility of a noble gas in contact with silicate melt.
[15] The range of concentrations in which Henry's law applies becomes narrower the more the system diverges from ideal behavior.
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.
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).
In mathematical terms: Raoult's law can also be related to non-gas solutes.
