Biochemists usually discuss acids and bases in terms of their ability to donate and accept protons; that is, they use the Brønsted definition of acids and bases. A few concepts from general chemistry are important to help organize your thoughts about biochemical acids and bases:
Many living organisms (there are many exceptions among the microbes) can exist only in a relatively narrow range of pH values. Thus, vegetables are often preserved by pickling them in vinegar, a dilute solution of acetic acid in water. The low pH of the solution prevents many bacteria and molds from growing on the food. Similarly, it is a cliche of movie Westerns that desert springs whose water is alkaline (basic) are decorated with the skulls of cattle who were unfortunate enough to drink from them. Finally, individuals with chest injuries who are unable to breathe efficiently develop a metabolic acidosis, as their blood pH drops below normal due to the impaired elimination of CO2 (a weak acid) from the lungs.
Microorganisms capable of living in acidic environments expend a large amount of energy to keep protons from accumulating inside their membranes. These examples show the importance of controlling the pH of biological systems: Biochemical reactions, and therefore life, can exist only in a narrow, near-neutral pH range.
All physiological pH control relies ultimately on the behavior of weak acids and bases as buffers. A buffer is a combination of a weak acid and its salt or a weak base and its salt. The addition of an acid or a base to a buffered solution results in a lesser pH change than would occur if the acid were added to water alone.
This behavior is described quantitatively by the Henderson-Hesselbach equation,which can be derived from the definition of Ka: