Ionic Equilibria and pH

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:

H 2 O + H 2O <=> H 3O+ + 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:

H 2 O <=> H+ + OH-

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

> K w = [H+][OH-] = 1 x 10-14 M2

where [H+] is the hydrogen ion concentration and [OH-] is the hydroxide ion concentration. Because K w 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.

See also: Henderson-Hasselbalch Equation, The Structure and Properties of Water