Ionization of Water, Weak Acids, and Weak Bases: -Weak Acids and Bases Have Characteristic Dissociation Constants
Hydrochloric, sulfuric, and nitric acids, commonly called strong acids, are completely ionized in dilute aqueous solutions; the strong bases NaOH and KOH are also completely ionized. Of more interest to biochemists is the behavior of weak acids and bases—those not completely ionized when dissolved in water. These are common in biological systems and play important roles in metabolism and its regulation. The behavior of aqueous solutions of weak acids and bases is best understood if we first define some terms. Acids may be defined as proton donors and bases as proton acceptors. A proton donor and its correspon ding proton acceptor make up a conjugate acid-base pair (Fig. 2–16). Acetic acid (CH3COOH), a proton donor, and the acetate anion (CH3COO-), the corresponding proton acceptor, constitute a conjugate acid base pair, related by the reversible reaction
Each acid has a characteristic tendency to lose its proton in an aqueous solution. The stronger the acid, the greater its tendency to lose its proton. The tendency of any acid (HA) to lose a proton and form its conjugate base (A-) is defined by the equilibrium constant (Keq) for the reversible reaction
Equilibrium constants for ionization reactions are usu ally called ionization or dissociation constants, often designated Ka. The dissociation constants of some acids are given in Figure 2–16. Stronger acids, such as phosphoric and carbonic acids, have larger dissociation constants; weaker acids, such as monohydrogen phosphate (HPO42-), have smaller dissociation constants.
FIGURE 2–16 Conjugate acid-base pairs consist of a proton donor and a proton acceptor. Some compounds, such as acetic acid and ammonium ion, are monoprotic; they can give up only one proton. Others are diprotic (H2CO3 (carbonic acid) and glycine) or triprotic (H3PO4 (phosphoric acid)). The dissociation reactions for each pair are shown where they occur along a pH gradient. The equilibrium or dis sociation constant (Ka) and its negative logarithm, the pKa, are shown for each reaction.
Also included in Figure 2–16 are values of pKa, which is analogous to pH and is defined by the equation
The stronger the tendency to dissociate a proton, the stronger is the acid and the lower its pKa. As we shall now see, the pKa of any weak acid can be determined quite easily.
