Dalton’s Law of Partial Pressures

John Dalton formulated this law around 1802, based on his experiments with gases. It was a crucial step in the development of his atomic theory, as it suggested that the particles of different gases in a mixture act independently of one another. The law holds true for ideal gases and is a good approximation for real gases at low pressures and high temperatures, where intermolecular interactions are minimal. The partial pressure of a single gas ([latex]p_i[/latex]) in a mixture can be calculated using its mole fraction ([latex]x_i[/latex]) and the total pressure ([latex]P_{text{total}}[/latex]): [latex]p_i = x_i cdot P_{text{total}}[/latex]. This relationship is extremely useful in many practical scenarios. For instance, in scuba diving, the partial pressure of oxygen is critical; if it becomes too high at depth, it can cause oxygen toxicity. Divers use Dalton’s law to calculate the maximum safe operating depth for a given breathing gas mixture. In physiology, the exchange of oxygen and carbon dioxide in the alveoli of the lungs is driven by differences in their partial pressures between the air and the blood. The law also explains why collecting a gas over water requires a correction, as the total measured pressure includes the partial pressure of water vapor at that temperature, which must be subtracted to find the pressure of the collected gas alone.