KEY2CHEM

Comparing Strong and Weak Acids

Strong acids are acids that dissociate completely in water, while weak acids are those that dissociate only slightly. The exact extent of dissociation depends on the $$K_a$$ of the acid. Remember that $$K_a$$ is the acid dissociation constant, describing the dissociation of an acid $$\require{mhchem}\ce{HA}: \require{mhchem}\ce{HA(aq) + H2O(l) <=> H3O+(aq) + A^{-}(aq)}$$. The percent of the acid that dissociates (percent ionization) depends on the $$K_a$$ of the acid and the initial concentration. The stronger the acid, the larger percent ionization.

Since strong acids dissociate to a greater extent than weak acids, they generate a higher concentration of $$\require{mhchem}\ce{H3O+}$$, leading to a lower $$pH$$ when comparing equal intial acid concentrations.

For a weak acid solution and a strong acid solution with the same initial concentration, it takes much more base to neutralize the weak acid solution because the initial acid concentration is much larger. The weak acid solution contains a large amount of un-ionized acid molecules. Therefore, a weak acid solution resists changes in for a much greater amount of added base.

Example 1.

Which acid will require the largest volume of of $$0.1\text{ M } \require{mhchem}\ce{NaOH}$$ to be fully neutralized?

A$$0.1 \text{ M } \require{mhchem}\ce{HCl}\;\;\; K_a = 1.3 \times 10^6$$

B. $$0.1 \text{ M } \require{mhchem}\ce{HBr}\;\;\;\; K_a = 1.0 \times 10^9$$

C. $$0.1 \text{ M } \require{mhchem}\ce{HF}\;\;\;\; K_a = 7.2 \times 10^{-4}$$

Solution

C. $$0.1 \text{ M } \require{mhchem}\ce{HF}\;\; K_a = 7.2 \times 10^{-4}$$

All of these acids have the same initial concentration, so the weakest acid (the one with the smallest $$K_a$$) will require the largest amount of base to be fully neutralized.

Example 2.

Which acid has the greatest percent ionization?

A. $$0.1 \text{ M } \require{mhchem}\ce{HF}\; K_a = 7.2 \times 10^{-4}$$

B. $$0.1 \text{ M } \require{mhchem}\ce{HClO2}\; K_a = 1.1 \times 10^{-2}$$

C. $$0.1 \text{ M } \require{mhchem}\ce{HCOOH}\; K_a = 1.8 \times 10^{-4}$$

Solution

B. $$0.1 \text{ M } \require{mhchem}\ce{HClO2}\; K_a = 1.1 \times 10^{-2}$$

All of these acids have the same initial concentration, so the strongest acid (the one with the largest $$K_a$$ value) will have the largest percent ionization.

Example 3.

Which solution will have the lowest pH?

A. $$0.1 \text{ M } \require{mhchem}\ce{HF}\; K_a = 7.2 \times 10^{-4}$$

B. $$0.1 \text{ M } \require{mhchem}\ce{HIO3}\; K_a = 1.6 \times 10^{-1}$$

C. $$0.1 \text{ M } \require{mhchem}\ce{HBrO}\; K_a = 2.3 \times 10^{-9}$$

Solution

B. $$0.1 \text{ M } \require{mhchem}\ce{HIO3}\; K_a = 1.6 \times 10^{-1}$$

All of these acids have the same initial concentration, so the strongest acid (the one with the largest $$K_a$$ value) will have the lowest pH (greatest [$$\require{mhchem}\ce{H3O+}$$]).