General Chemistry/Buffer Systems

Introduction
Buffer systems are systems in which there is a significant (and nearly equivalent) amount of a weak acid and its conjugate base—or a weak base and its conjugate acid—present in solution. This coupling provides a resistance to change in the solution's pH. When strong acid is added, it is neutralized by the conjugate base. When strong base is added, it is neutralized by the weak acid. However, too much acid or base will exceed the buffer's capacity, resulting in significant pH changes.

Buffers are useful when a solution must maintain a specific pH. For example, blood is a buffer system because the life processes in a human only function within a specific pH range of 7.35 to 7.45. When, for example, lactic acid is released by the muscles during exercise, buffers within the blood neutralize it to maintain a healthy pH.

Making a Buffer
Once again, let's consider an arbitrary weak acid, HA, which is present in a solution. If we introduce a salt of the acid's conjugate base, say NaA (which will provide the A- ion), we now have a buffer solution. Ideally, the buffer would contain equal amounts of the weak acid and conjugate base.

Instead of adding NaA, what if a strong base were added, such as NaOH? In that case, the hydroxide ions would neutralize the weak acid and create water and A- ions. If the solution contained only A- ions, then a strong acid like HCl were added, they would neutralize and create HA.

As you can see, there are three ways to create a buffer:

All six of the combinations will create equal amounts of a weak acid and its conjugate base, or a weak base and its conjugate acid.

Buffers and pH
To determine the pH of a buffer system, you must know the acid's dissociation constant. This value, $$K_a$$ (or $$K_b$$ for a base) determines the strength of an acid (or base). It is explored more thoroughly in the Equilibrium unit, but for now it suffices to say that this value is simply a measure of strength for acids and bases. The dissociation constants for acids and bases are determined experimentally.

The Henderson-Hasselbalch equation allows the calculation of a buffer's pH. It is:

$$\hbox{pH} = \hbox{pK}_a + \log \frac{[\hbox{A}^-]}{[\hbox{HA}]}$$

For a buffer created from a base, the equation is:

$$\hbox{pH} = \hbox{pK}_b + \log \frac{[\hbox{B}]}{[\hbox{HB}^+]}$$

Using these equations requires determining the ratio of base to acid in the solution.