William Henry, an English chemist created the gas law from the studies that he conducted in the initial moments of the 19^{th} century. He arrived at the conclusion that a gas’s partial pressure in its phase is proportion to the amount of dissolved gas. The mathematical expression of Henry’s law is:

*p ***= ***k*_{H}**⋅***c*

Above the solution, the solute’s partial pressure is determined as *p*

The solutes concentration in the solution is represented by *c*

The Henry’s Law Constant is *k*_{H}

The various values of ** k_{H }**are:

- 2 L · atm/mol: O
_{2}(oxygen) - 4 L · atm/mol: Carbon Dioxide (CO
_{2}) - 1 L · atm/mol: Hydrogen (H
_{2})

This is all happening at a temperature of 298 K.

## Henry’s Law Constant and its dependence on Temperature

The Henry’s law of constant will change with the change in the temperature of the system. The temperature’s effect on the constant is assessed by multiple equations. Some of the examples of the equations that are used to represent it are Van’t Hoff equation:

The Henry’s Law Constant is *k*_{H}

A constant that is measured in Kelvin’s is **C**

Temperature is **T (Kelvin)**

**298K is the standard state temperature or T**

^{ }

^{Θ}**The equations mentioned above are used when a given gas has no other experimentally derived formula. **

The phenomenon of decompression sickening is an instance, where Henry’s law comes into play. When a person dives into the water and swims too deep, suffers the changes in compression that affect the oxygen and nitrogen content in the blood. A similar instance can be seen with soft drinks that have been carbonated by the use of carbon dioxide. The carbon dioxide is filled in the space above the drink, which has a higher pressure than the atmospheric pressure. The dissolved carbon dioxide undergoes degassing as the cap is opened and the partial pressure of carbon dioxide is lowered.

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