Kaplan MCAT General Chemistry Review

In region IV, the acid has neutralized approximately half of the HCO₃^-, and now H₂CO₃ and HCO₃^- are in roughly equal concentrations. This flat region is the second buffer region of the titration curve, corresponding to the pK_a of H₂CO₃ (K_a = 4.3 × 10^-7 implies pK_a = 6.37). In region V, the equivalence point for the entire titration is reached, as all of the HCO₃^- is finally converted to H₂CO₃. Again, a rapid change in pH is observed near the equivalence point as acid is added.

The titrations of the acidic and basic amino acids (which have acidic or basic side chains, respectively) will show curves similar to the one shown in Figure 10.3. But rather than two equivalence points, there will in fact be three: one corresponding to the titration of the carboxylic acid and a second corresponding to the titration of the amino acid, both of which are attached to the central carbon, and a third corresponding either to the acidic or basic side chain.

BUFFERS

A buffer solution consists of a mixture of a weak acid and its salt (which consists of its conjugate base and a cation) or a mixture of a weak base and its salt (which consists of its conjugate acid and an anion). Two examples of buffers that are common in the laboratory and commonly tested on the MCAT are a solution of acetic acid (CH₃COOH) and its salt, sodium acetate (CH₃COO^- Na⁺), and a solution of ammonia (NH₃) and its salt, ammonium chloride (NH₄⁺Cl^-). The acetic acid/ sodium acetate solution is an acid buffer, and the ammonium chloride/ammonia solution is a base buffer. Buffer solutions have the useful property of resisting changes in pH when small amounts of strong acid or base are added. Consider a buffer solution of acetic acid and sodium acetate (the sodium ion has not been included because it is not involved in the acid-base reaction):

CH₃COOH (aq) + H₂O (l) H₃O⁺ (aq) + CH₃COO^- (aq)

When a small amount of strong base, such as NaOH, is added to the buffer, the OH^- ions from the NaOH react with the H⁺ ions present in the solution; subsequently, more acetic acid dissociates (the system shifts to the right), restoring the [H⁺]. The weak acid component of the buffer acts to neutralize the strong base that has been added. The resulting increase in the concentration of the acetate ion (the conjugate base) does not yield as large an increase in hydroxide ions as the unbuffered strong base would. Thus, the addition of the strong base does not result in a significant increase in [OH^-] and does not appreciably change pH. Likewise, when a small amount of HCl is added to the buffer, H⁺ ions from the HCl react with the acetate ions to form acetic acid. Acetic acid is weaker than the added hydrochloric acid (which has been neutralized by the acetate ions), so the increased concentration of acetic acid does not significantly contribute to the hydrogen ion concentration in the solution. Because the buffer maintains [H⁺] at relatively constant values, the pH of the solution is relatively unchanged.

In the human body, one of the most important buffers is the H₂CO₃/HCO₃^- conjugate pair in the plasma component of the blood. CO₂ (g), one of the waste products of cellular respiration, has low solubility in aqueous solutions. The majority of the CO₂ transported from the peripheral tissue to the lungs (where it will be exhaled out) is dissolved in the plasma in a “disguised” form. CO₂ (g) and water react in the following manner:

CO₂ (g) + H₂O (l) H₂CO₃ (aq) H⁺ (aq) + HCO₃^- (aq)

Carbonic acid (H₂CO₃) and its conjugate base, bicarbonate (HCO₃^-), form a weak acid buffer for maintaining the pH of the blood within a fairly narrow physiological range. The most important point to notice about this system for pH homeostasis is its direct connection to the respiratory system. In conditions of metabolic acidosis (excess of plasma H⁺), for example, the respiratory rate (breathing rate) will increase in order to “blow off” a greater amount of carbon dioxide gas; this causes the system to shift from the right to the left, thereby restoring the normal physiological pCO₂ and in doing so, reducing the [H⁺] and buffering against dramatic and dangerous changes to the blood pH. One of the more interesting topics to ponder about the blood buffer system is why a weak acid buffer system was selected (evolutionarily speaking) as a primary mechanism for human blood pH homeostasis at around pH 7.4 (weakly basic). Buffers have a definite and narrow range of optimal buffering capability (pK_a ±1), and pH 7.4 is actually slightly above the outer limit of buffering capability for the carbonic acid/bicarbonate system (pK_a = 6.1). It’s an interesting question—and one to which there is an answer—but we don’t want to ruin all the surprises of medical school, so we will leave it unanswered for the time being.