Common Ion Effect in the Ionization of Acids and Bases

The Common Ion Effect in Acid-Base Ionization

When a weak acid or weak base is placed in water, it establishes an equilibrium between the undissociated molecule and its ions. The position of this equilibrium is not fixed — it shifts in response to changes in concentration, just as Le Chatelier's principle predicts. The common ion effect is the name given to the shift that occurs when one of the ions already present in the equilibrium is added from an external source.

The Acetic Acid Example

Consider acetic acid, a weak monoprotic acid. Its dissociation equilibrium in water is:

or, using the shorthand notation HAc for acetic acid and Ac⁻ for acetate ion:

The acid dissociation constant for this equilibrium is:

At a given temperature, $K_a$ is a fixed number. This means that if we change the concentration of any species in the expression, the other concentrations must adjust to keep the ratio constant.

What happens when acetate ions are added? Suppose we add sodium acetate (NaAc) to the solution. Sodium acetate dissociates completely, providing a large concentration of Ac⁻ ions. The equilibrium quotient $[\mathrm{H^+}][\mathrm{Ac^-}]/[\mathrm{HAc}]$ temporarily exceeds $K_a$. To restore equilibrium, the system shifts to the left — more HAc is formed, consuming H⁺ and Ac⁻. The net result is a decrease in $[\mathrm{H^+}]$.

What happens when H⁺ ions are added? If we add a strong acid like HCl, the concentration of H⁺ rises sharply. Again, the quotient exceeds $K_a$, and the equilibrium shifts left, converting H⁺ and Ac⁻ into undissociated HAc. This reduces the concentration of free H⁺, though not back to its original value.

Why This Matters

The common ion effect is the fundamental principle behind buffer solutions. A buffer resists changes in pH because it contains both a weak acid and its conjugate base (or a weak base and its conjugate acid). When an external acid or base is added, the common ion effect shifts the equilibrium to consume most of the added H⁺ or OH⁻, keeping the pH nearly constant.

A Quantitative Look

Let's see how the common ion effect changes the pH of a weak acid solution. For a 0.10 M acetic acid solution alone, the equilibrium calculation gives:

Now suppose we add enough sodium acetate to make the solution 0.10 M in Ac⁻ as well. Let $x$ be the concentration of H⁺ at equilibrium. The initial concentrations (after adding acetate but before any shift) are:

Substitute into the $K_a$ expression:

Because $K_a$ is very small, $x$ will be much smaller than 0.10. We can approximate $0.10 + x \approx 0.10$ and $0.10 - x \approx 0.10$:

So $x = 1.8 \times 10^{-5}$ M, and $\mathrm{pH} = -\log(1.8 \times 10^{-5}) = 4.74$.

The General Principle

For any weak acid HA:

Adding A⁻ (from a salt like NaA) shifts the equilibrium left, decreasing [H⁺]. Adding H⁺ (from a strong acid) also shifts the equilibrium left, decreasing [A⁻] and increasing [HA]. …