• BME 200

# Henry’s Law

Henry’s law says that the concentration of a dissolved gas depends on its partial pressure, temperature, and solubility. Imagine that a beaker half full of water has been sitting in a vacuum chamber overnight so that no gas is dissolved in the water (Figure 1a). We take the beaker out and let it sit in air at room temperature (20 $$^\circ$$C) and pressure (760 mmHg) until it equilibrates, as shown in Figure 1b. Once the system reaches equilibrium, the partial pressure of oxygen will be the same in the air and the water, but the concentration will be different. The water molecules take up space, so there cannot be as many oxygen molecules per unit volume in the water even though the partial pressures are equal. At normal temperature and pressure (20 $$^\circ$$C and 760 mmHg) the concentration of oxygen in air is 8.61 mol/m$$^3$$ and in water at is 0.21 mol/m$$^3$$. Figure 1. (a) Water in a beaker placed in a vacuum has no oxygen. (b) At equilibrium with room air the partial pressure in the water and in the air above the water is the same. However, the concentration of oxygen in the two phases is different.

Henry’s law is used to relate the concentration of a gas to its partial pressure. The affinity that a gas molecule has for a liquid or solid phase is called its solubility coefficient, or $$\alpha$$. Henry’s law is

$$P_A = \frac{C_A}{H}$$

where $$P_A$$ is the partial pressure of gas $$A$$, $$C_A$$ is the concentration of gas $$A$$, and $$H$$ is solubility coefficient of dissolved gas $$A$$. At STP $$H=1.005 \times 10^{-11}~\textrm{mol}~\textrm{cm}^{-3}~\textrm{Pa}^{-1}.$$

Physiologists refer to the partial pressure of a gas in solution as its oxygen tension.

Last updated:
August 25, 2018