Acids, Bases, and pH

STRONG AND WEAK ACIDS

Ionization

There are two classes of acids: strong acids and weak acids. Note that strong acids and bases are rarely a concern in biochemistry (as they are in general chemistry). Strong acids ionize (break into ions) completely when dissolved in water. In other words, a strong acid is a good proton donor. For example, the strong acid HCl gives up all its protons to water:

HCl + H2O Cl– + H3O+

There are only a few strong acids, which are all listed below. HCl (hydrochloric acid, found in stomach digestive juices) is a good example of a strong acid.

Because strong acids completely ionize, the [H⁺] of a solution made with a strong acid is easily figured out, since it is equal to the molarity of the solution. Thus, a 0.1 M solution of HCl has a [H⁺] of 0.1. A 0.029 M solution of HCl has a [H⁺] of 0.029 and so on. (Remember that this only holds true as long as the molarity of your solution is somewhat greater than 10^–7 M. Otherwise, for solutions so dilute that the H⁺ normally present in pure water is comparable or larger, you will need to take the [H⁺] of pure water into consideration.)

Similar to the case with strong acids, there are just a few strong bases that ionize completely. For example, when sodium hydroxide ionizes in water, the resulting hydroxide ion readily accepts a proton:

NaOH Na+ + OH–

The strong bases are listed in the table below.

Most acids of biological origin are weak. In part 1 of this tutorial, we saw that in pure water, a very small percentage of the H₂O molecules dissociate to form H⁺ and OH^– ions. Most of the molecules in pure water remain intact as H₂O. The same can be said of weak acids. When weak acids are dissolved in water, only some of the acid molecules actually dissociate (break apart), while most others remain intact.

The percentage of the molecules that break apart depends on the weak acid. Some weak acids dissociate very minimally, and as a consequence such an acid has little acid strength. Other weak acids dissociate more, and as a consequence have more acid strength.

The degree of dissociation of a weak acid in water is described by the acid dissociation constant, K_a.

For the dissociation of acetic acid (the acid in vinegar) in water,

CH3COOH + H2O CH3COO– + H3O+

the equilibrium (or acid dissociation) constant expression for the reaction is given by

(Remember that by convention, [H₂O] is left out of the expression because the reaction is taking place in water). The “a” subscript on K_a is a reminder that this number represents an acid dissociation constant. The larger the K_a, the more the acid dissociates (or reacts with water) to form H₃O⁺ ions, and the stronger the acid.

Note how K_a values for weak acids tend to be very small numbers. Because these numbers with exponents tend to be cumbersome, K_a values are usually converted to pK_a values, much in the same way as [H⁺] values are converted to pH.

For example, the K_a of acetic acid is 1.74 x 10^–5.

The pK_a for acetic acid is therefore pK_a = –log (1.74 x 10^–5) = 4.76.

Below is a table listing some common weak acids and their K_a and pK_a values.

Note that the smaller the K_a, the larger the pK_a. Thus, stronger acids are represented by larger K_a values, but smaller pK_a values.

Also, whether an acid is strong of weak can be readily identified by either its K_a or pK_a:

Hydronium acid (H₃O⁺), which is the protonated form of water, is the dividing line between a strong or weak acid. In other words, a strong acid is defined as one that ionizes to a larger degree than the H₃O⁺ acid form of water does (K_a = 1, pK_a = 0). Another way to put this is that if you put a strong acid in water, most of it ionizes, while for a weak acid most stays in its conjugate acid (protonated) form.