Henry's Law

In chemistry, Henry's law is one of the gas laws formulated by the English chemist , 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:

In other words, 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.

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 , 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 gas phase into the denominator ("aq/gas"). This results in the Henry's law solubility constant ${\displaystyle H}$. Its value increases with increased solubility. Alternatively, numerator and denominator can be switched ("gas/aq"), which results in the Henry's law volatility constant ${\displaystyle K_{\rm {H}}}$. The value of ${\displaystyle K_{\rm {H}}}$ 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 (${\displaystyle c_{\rm {a}}}$), molality (${\displaystyle b}$), and molar mixing ratio (${\displaystyle x}$). For the gas phase, molar concentration (${\displaystyle c_{\rm {g}}}$) and partial pressure (${\displaystyle p}$) are often used. It is not possible to use the gas-phase mixing ratio (${\displaystyle y}$) because at a given gas-phase mixing ratio, the aqueous-phase concentration ${\displaystyle c_{\rm {a}}}$ depends on the total pressure and thus the ratio ${\displaystyle y/c_{\rm {a}}}$ 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, ${\displaystyle H^{cp}}$ refers to the Henry solubility defined as ${\displaystyle c/p}$.

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