pH, Carbon Dioxide, and Alkalinity

Abstract

The pH or negative logarithm of the hydrogen ion concentration is a master variable in water quality because the hydrogen ion influences many reactions. Because dissolved carbon dioxide is acidic, rainwater that is saturated with this gas is naturally acidic—usually about pH 5.6. Limestone, calcium silicate, and feldspars in soils and other geological formations dissolve through the action of carbon dioxide to increase the concentration of bicarbonate in water and raise the pH. The total concentration of titratable bases—usually bicarbonate and carbonate—expressed in milligrams per liter of calcium carbonate is the total alkalinity. Total alkalinity typically is less than 50 mg/L in waters of humid areas with highly leached soils, but it is greater where soils are more fertile, limestone formations are present, or the climate is arid. Alkalinity increases the availability of inorganic carbon for photosynthesis, because most phytoplankton species and many higher aquatic plants can obtain carbon from bicarbonate. Water bodies with moderate to high alkalinity are well-buffered against wide daily swings in pH resulting from net removal of carbon dioxide by photosynthesis during daytime and return of carbon dioxide to the water by respiratory process at night when there is no photosynthesis. The optimum pH range for most aquatic organisms is 6.5–8.5, and the acid and alkaline death points are around pH 4 and pH 11, respectively. Fish and other aquatic animals avoid high carbon dioxide concentration, but 20 mg/L or more can be tolerated if there is plenty of dissolved oxygen.