Water, as a polar molecule, induces an accumulation of electron density (dipole moment) at one end of non-polar gas molecules such as oxygen (O₂ ) and carbon dioxide (CO₂ ). In animation, observe a polar water molecule approaching a nonpolar O₂ molecule. The electron cloud of O₂ is normally distributed symmetrically between the bonded O₂ atoms. When the negative end of the H₂O molecule approaches the oxygen molecule, the electron cloud of the O₂ moves away to reduce the negative-to-negative repulsion. A dipole (a molecule with positive and negative charges separated by a distance) results in the nonpolar O₂ molecule and causes O₂ and H₂O to become weakly attracted to each other. This intermolecular attraction between the oppositely charged poles of nearby molecules is termed a dipole- dipole force. The creation of these forces explains the mechanism by which gases dissolve in water.

2. The effects of pressure on oxygen solubility

Because dipole- induced dipole forces are very weak, the quantity of nonpolar gases (such as O₂) that will dissolve in a given volume of water is strongly affected by temperature and pressure. Henry's Law describes the effect of pressure on the solubility of a gas in a liquid. The law states that the amount of gas that dissolves in a given volume of solvent at a specified temperature (usually 25°C for water) is proportional to the partial pressure of the gas above the liquid. When gas under pressure contacts a liquid, the pressure tends to force gas molecules into solution. At a given pressure the number of gas molecules that will enter into solution rises until equilibrium is reached. By definition, at equilibrium, the number of gas molecules entering and leaving the solutions is balanced and the concentration of the gas in solution remains constant. If the partial pressure of a gas increases, more gas enters into solution. If partial pressure drops, gas comes out of solution and reaches a new equilibrium. This can be illustrated by opening a can or bottle of soda pop. (See note under 3.)

At sea level total atmospheric pressure is 760 mm Hg. This means that the gravity-induced weight of the atmosphere generates enough force to move a sufficient volume of mercury (Hg) 760 mm up a tube. At sea level approximately 20.8 percent of this air is oxygen gas (O₂). The partial pressure of oxygen at sea level is 158 mm Hg (760 mm Hg x 0.208 = 158.08 mm Hg). Oxygen has a Henry's Law constant of 1.7 x 10-6 molal/mm Hg when dissolved in water at 25°C. Molality of O₂ = (1.7 x 10-6 molal/mm Hg) (158.08 ) = 2.687 x 10-4 m.

From the above value the number of milligrams per liter of oxygen that will dissolve in 25°C water can be calculated. 2.687 x 10-4 moles/kg x 32g/mole x 1000 mg/g = 8.6 mg/liter

3. The effects of temperature on solubility - Le Chatelier's Principle.

Pour cold water into a glass. After the animation you can observe the oxygen bubbles that formed inside the glass of water that had been poured cold and warmed up over time. Next, open two cans of soda pop, one warm and one cold. It is easily observed that more gas is released if the can is warm than when it is cold. Both of demonstrations illustrate that the temperature of a solvent (recall that water is the "universal solvent") affects the solubility of gases.

Gases that dissolve in solvents usually release heat as they dissolve in an exothermic process. This process continues until saturation is reached. At this point gas will still dissolve, but will be balanced by the gas that leaves solution. If heat is added to a solution, gas is released in this endothermic reaction:

At equilibrium, as many molecules come out of solution as dissolve in a given time period.

Le Chatelier's principle states that a change in any of the factors determining equilibrium will cause the system to adjust in order to reduce or counteract the effect of the change. Le Chatelier's principle predicts that the solubility of a gas will increase as a system loses heat and decrease as it gains heat.

4. O₂ supersaturation

Due to the effects of hydrostatic pressure on gases in solution, water can become supersaturated (exceed 100% saturation) with oxygen and other gases. The attractive forces that hold excess oxygen in solution are similar to the dipole-induced dipole forces discussed earlier, but a smaller number of water molecules are available to induce dipoles in oxygen molecules. This leads to weaker attraction of oxygen molecules when water is supersaturated with oxygen.