Biochemistry For Dummies. John T. Moore

      Acting as either an acid or a base

Some substances can’t make up their minds about what they are; they can act as either an acid or a base. Chemists classify these substances as amphiprotic or amphoteric substances. For example, the bicarbonate ion can act as either an acid or a base:

       

       

      Biochemically important molecules may also exhibit amphiprotic behavior. Amino acids contain both a basic amine group and an acidic carboxyl group. Therefore, they can act as either acids or bases. For example, glycine may undergo the following reactions:

       

       

      In fact, amino acids may undergo proton transfer from the carboxyl end to the amine end, forming an overall neutral species that has a positive and negative end. Species such as these are called zwitterions (not to be confused with twitterions — people who tweet until their thumbs fall off):

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:

       Respiratory acidosis is the result of many diseases that impair respiration, including pneumonia, emphysema, and asthma. These diseases are marked by inefficient expulsion of carbon dioxide, leading to an increase in the concentration of carbonic acid, H2CO3.

       Metabolic acidosis is due to a decrease in the concentration of (the bicarbonate ion). This decrease may be the result of certain kidney diseases, uncontrolled diabetes, and cases of vomiting involving nonacid fluids. Poisoning by an acid salt may also lead to metabolic acidosis.

       Respiratory alkalosis may result from hyperventilation because this excessive removal of carbon dioxide can lead to a decrease in the H2CO3 concentration. Immediate treatment includes breathing into a paper bag, which increases the carbon dioxide concentration in the inhaled air and, therefore, in the blood.

       Metabolic alkalosis may result from excessive vomiting of stomach acid.

      

To resist these pH problems, the blood has a number of buffer systems — systems that resist a change in pH by reacting with either added acids or bases. In general, buffers may be amphiprotic substances or mixtures of weak acids and weak bases. In the body, these include several proteins in blood plasma and the bicarbonate buffer system.

      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