Question

The \(\mathrm{pH}\) of a \(0.0642 \mathrm{M}\) solution of a monoprotic acid is \(3.86 .\) Is this a strong acid?

Short answer

No, this is not a strong acid.

Step-by-step solution

Calculate the concentration of \(H^{+}\)

Determine the concentration of \(H^{+}\) in the solution using the expression \( [H^{+}]=10^{-\mathrm{pH}}\). Here, \(\mathrm{pH}=3.86\). So, \( [H^{+}]=10^{-3.86}\). The calculation yields \( [H^{+}]\) equals approximately \(1.37 \times 10^{-4} \mathrm{M}\).

Compare \(HA\) and \(H^{+}\)

Compare the concentration of the original acid (\(HA=0.0642 \mathrm{M}\)) with that of the H+ ions (\(H^{+}=1.37 \times 10^{-4} \mathrm{M}\)). It is evident that \(HA >> H^{+}\). This means that only a small fraction of the original acid has ionized into \(H^{+}\) ions.

Determine acid strength

Based on the comparison in Step 2, it can be concluded that this is not a strong acid. Strong acids ionize completely and would have a \(H^{+}\) concentration that is equal to the starting concentration of the acid.