Sodium Sulfate vs Sodium Sulfite

I recently took a test and one of the questions was:

Which procedures will allow a student to differentiate between solid sodium sulfate and solid sodium sulfite?

I. Make solutions of each and look for a precipitate when added to $$0.10$$ M barium nitrate.

II. Add crystal of each to $$0.10$$ M HCl and watch for bubbles.

III. Make solutions of each and test with a pH indicator.

Why is III also correct, if sodium sulfate and sodium sulfite are both basic and would turn pH indicator the same color?

The last statement

…sodium sulfate and sodium sulfite are both basic and would turn pH indicator the same color

is true only for $$\ce{Na2SO4}$$ as it's formed by both strong base and strong acid and won't noticeably affect pH. $$\ce{Na2SO3}$$, on the other hand, undergoes hydrolysis:

\begin{align} \ce{Na2SO3 + H2O &<=> NaOH + NaHSO3} \\ \ce{2Na+ + SO3^2- + H2O &<=> Na+ + OH- + Na+ + HSO3-} \\ \ce{SO3^2- + H2O &<=> HSO3- + OH-} \end{align}

resulting in elevated pH (basic solution) up to 9 (according to Wikipedia).

• But sulfate is a weak base too, right? (HSO4- is a weak acid) – Cyclopropane Apr 16 '19 at 16:15

The questions asks you to differentiate Na2SO4 and Na2SO3. There are two main concepts involved here:

H2SO3 is a very weak acid (better to write SO2 (aq)) and H2SO4 is an extremely strong acid. Certainly, HCl would not be able to displace H2SO4. Another point which you should remember is that salts of strong acids and strong bases are pH neutral. However, once there is a disbalance in the strength of acids or bases, their corresponding salts solutions are either weakly basic or acidic. Salts or weak bases and weak acids are quite complicated in terms of pH calculations.

A solution of Na2SO3, formed from a strong base and weak acid, would show alkaline behavior.