Buffers and pH Control

      A buffer solution contains the conjugate acid-base pair of any weak acid or base in relative proportions to resist pH change when small amounts of either an acid or a base are added. Therefore, buffers control the pH of the solution. Buffer solutions are important in most biological systems. Many biological processes proceed effectively only within a limited pH range. The presence of buffer systems keeps the pH within this limited range.

      Identifying common physiological buffers

      In the human body, the pH of various body fluids is important. The pH of blood is 7.4, the pH of stomach acid is , and the pH in the intestinal tract is . If the pH of blood is more than 0.2 pH units lower than normal, a condition known as acidosis results; a corresponding increase in pH of about the same magnitude is alkalosis. Acidosis and alkalosis, which may lead to serious health problems, each have two general causes:

      The bicarbonate buffer system is the main extracellular buffer system. This system also provides a means of eliminating carbon dioxide. The dissolution of carbon dioxide in aqueous systems sets up the following equilibrium:

      The presence of the conjugate acid-base pair (H₂CO₃ and ) means that this is a buffer system. The conjugate acid-base ratio is about 20:1 at a pH of 7.4 in the bloodstream. This buffer system is coupled with the following equilibrium (instrumental in the removal of carbon dioxide in the lungs):

      The second ionization of phosphoric acid, K_a2, is the primary intracellular buffer system. The pH of this conjugate acid-base pair ( and ) is 7.21 for a solution with equal concentrations of these two species.

      Calculating a buffer’s pH

      To determine a buffer’s pH, you can use a K_a or K_b calculation, as we discuss in the section “Swapping hydrogens between acids and bases,” earlier in this chapter, or you can use the Henderson-Hasselbalch equation, which is a shortcut.

      The Henderson-Hasselbalch equation takes two forms:

      and

      The terms in either form are the same as those we define earlier in the chapter. For example, suppose you want to calculate the pH of a buffer composed of 0.15 M pyruvic acid and 0.25 M sodium pyruvate. Referring back to Table 2-2, you see that the K_a of pyruvic acid is .

      The pK_a would be 2.50. Therefore:

      

The greater the values of [CA] and [CB], the greater the buffer capacity of the solution. The buffer capacity indicates how much acid or base may be added to a buffer before the buffer ceases to function as a buffer. A buffer in which the would have a much higher buffer capacity for adding either acids or bases than a buffer in which the . If there were a buffer in which and its , the buffer would have a higher buffer capacity for additions of a base than for additions of an acid because the buffer contains more acid than base. For the buffer to be as flexible as possible, the concentrations of the conjugate acid-base pair should be as close to equal as possible

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