Diffusion is due to the random motion of molecules, and it results in a net movement of molecules from areas of higher concentration to areas of lower concentration. In other words, diffusion evens out the distribution of molecules. Diffusion of respiratory gases occurs either in water or in air. Concentrations of gases in water or air vary with pressure because gases are compressible, a relationship described in physics by Boyle’s Law. For example, there are twice as many gas molecules in a liter of gas at 2 atmospheres of pressure as there are in a liter of gas at 1 atmosphere of pressure. And if that gas is in contact with a liquid such as water, twice as many gas molecules will enter into solution when the gas pressure is 2 atmospheres than when it is 1 atmosphere.

The concentrations of different gases in a mixture are described as the partial pressures of those gases. To calculate the partial pressure of a gas such as oxygen in a mixture of gases such as air, we have to know the total pressure. The total pressure of air inhaled by air-breathing animals is the atmospheric pressure. At sea level, atmospheric pressure is about 760 millimeters of mercury (mm Hg), depending on the weather. Because dry air is 20.9 percent O₂, the partial pressure of oxygen (P_O2) at sea level is 20.9 percent of 760 mm Hg, or about 159 mm Hg. If two gas mixtures are separated by a membrane permeable to O₂, O₂ will diffuse from the mixture where its partial pressure is higher to the mixture where its partial pressure is lower.

To calculate the concentration of O₂ in the environment of a water breather, we have to know two things: the partial pressure of O₂ in the air in contact with the water, and the solubility of O₂ in water. The amount of a gas that dissolves in a liquid depends both on its partial pressure in the gas phase in contact with the liquid and on its solubility in that liquid. The diffusion of a gas between the gas phase and the liquid phase is a function of its partial pressures in those two phases; the gas diffuses from the phase with the higher partial pressure to the phase with the lower partial pressure until equilibrium is reached—the point at which the partial pressures in the two phases are equal. However, the amount of the gas that can be contained in the liquid depends on the solubility of that gas in that liquid. Furthermore, the solubility of a gas in a particular liquid can vary widely depending on conditions. What follows is a practical illustration of these facts.

Solubility of a gas in a liquid, such as oxygen in water, is a function of temperature—solubility is higher at low temperatures. Think of opening a bottle of warm soda in comparison to a bottle of cold soda. There is more gas out of solution—more fizz—in the warm soda. So if we have similar containers of water in equilibrium with air, but at different temperatures, the concentrations of oxygen in these water containers will be different (less O₂ in the warmer one), but the partial pressures of oxygen will be the same. Thus for water-breathing animals, the warmer the water is, the less O₂ there is per liter of water. The important point is that for a gas in solution, its concentration is not the same as its partial pressure, but partial pressures are what drive diffusion. Thus in our continuing discussions of respiratory gas exchange, we will always use partial pressure rather than concentration when referring to the diffusion of respiratory gases.