The bicarbonate buffering system and titration curves

The pH of blood in humans is around $7.4$‍. A rise of pH above $7.45$‍ leads to the condition of alkalosis that disrupt enzymes, causing muscle spasms and respiratory paralysis. Likewise, if physiological pH drops below $7.35$‍, it leads to acidosis that causes depression of the central nervous system. Several factors, including exercise, diet and changes in respiratory patterns, alter physiological pH. The body responds to these changes through the action of buffers that resist the alteration of pH. The bicarbonate buffering system in human blood, for example, maintains the pH around $7.4$‍ and is composed of carbonic acid ${\text{(H}}_2{\text{CO}}_3)$‍ and bicarbonate ion ${\text{(HCO}}_3^{-}\text{)}$‍.The buffer is formed when carbon dioxide dissolves into blood, forming carbonic acid. Carbonic acid and water then react to form hydronium ion ${\text{(H}}_3{\text{O}}^{+})$‍ and the bicarbonate conjugate base in solution.

Buffers can also be created in the laboratory by partially titrating either a weak acid with a strong base, or a weak base with a strong acid. This is visualized in a titration curve (Figure 1). After a sharp increase at the beginning, the pH during the course of the titration increases gradually due to the buffering capabilities of the solution. This continues until the base overcomes the buffers capacity after the equivalence point, and all of the initial acid has been converted to its conjugate base with the addition of the strong base.

Figure 1. The titration of a weak acid with strong base. The titration curve is a graph of the volume of titrant, or in our case the volume of strong base, plotted against the pH of the solution.

Attribution: J.A. Freyre, CC-BY-SA 2.5

The addition of an acid or a base to a buffered solution creates a smaller pH change than would occur if the acid or base were added to water alone. This behavior is described quantitatively by the Henderson-Hasselbalch equation (Figure 2), which can be used to deduce that pH of a buffer solution is equal to the ${\text{pK}}_a$‍ of the acid when equal molar concentrations of acid and its conjugate base is present. This point is also known as the half-equivalence point.

Figure 2. The Henderson-Hasselbalch equation