Ionic Equilibria - The common weak acids and bases found in biological systems do not completely ionize at physiological pH. Thus there is a measurable equilibrium between the weak acid and its conjugate base (the substance that can accept a proton to re-form the acid) or between the weak base and it conjugate acid (the stubstance that can give up a proton to re-form the base). Table 2.6 lists examples of weak acids and their conjugate bases.

Although water is essentially a neutral molecule, it does ionize slightly as follows:

H2O + H2O <=> H3O⁺ + OH^-

Thus, water acts as an acid (proton donor) and a base (proton acceptor) in what is called the autoionization of water. The autoionization of water can be simplified as follows:

H2O <=> H⁺ + OH^-

The equilibrium for the autoionization of water is given by Kw, the ion product constant of water. At 25C,

Kw = [H⁺][OH^-] = 1 x 10^-14 M²

where [H⁺] is the hydrogen ion concentration and [OH^-] is the hydroxide ion concentration. Because Kw is a constant, [H⁺] and [OH^-] can vary independently. If [H⁺] increases, then [OH^-] must decrease such that [H⁺][OH^-] = 1 x10^-14, and vice versa.

A solution is said to be neutral when [H⁺] = [OH^-]. At 25C, [H⁺] = [OH^-] = 1 x 10^-7M. Once [H⁺] is known, the pH of the solution can be calculated as follows:

pH = -log[H⁺]

For a neutral solution at 25C, pH = -log(1 x 10^-7) = 7. A solution with a pH <7 is said to be acidic, and one with a pH > 7 is said to be basic (see Figure 2.16). Most body fluids have pH values in the range 6.5 - 8.0, so this is often called the physiological pH range.