When I have a solution of $\ce{Na2SO4}$, the sodium sulfate should ionize: $$\ce{Na2SO4(aq) -> 2Na+(aq) + SO4^{2-}(aq)}$$

In my mind, the sodium ions should just float around, but because $\ce{HSO4-}$ is a weak acid:

$$\ce{HSO4- (aq) + H2O <-> H3O+ (aq) + SO4^2- (aq)}$$

The sulfate ion should be a weak conjugate base, and hence the sodium sulfate salt should be a weak base. This was my line of reasoning, until today my friends told me it is actually a neutral salt, which was confirmed by searching online. So why is it a neutral salt? (i.e. what's wrong with my reasoning?)

  • 3
    $\begingroup$ $\mathrm{HSO}_4^-$ can hardly be called weak. It is an acid of medium strength, stronger than carbonic or acetic. $\endgroup$ Sep 5 '15 at 9:51
  • $\begingroup$ Weak ions conjugate bases do not accept protons. $\endgroup$
    – Aditya Dev
    Sep 5 '15 at 10:06
  • $\begingroup$ Sulfate anions are weakly basic according to Bronsted theory - you're right, but it's called neutral salt according to Arrhenius theory in which it would need to contain OH- ions to be called basic. $\endgroup$
    – Mithoron
    Sep 5 '15 at 16:44
  • $\begingroup$ The sulfate ion is indeed a weak base, so weak that it hardly participates in the equilibrium reaction you have written. $\endgroup$
    – Zhe
    Jan 28 '17 at 2:15

For a solution of $\ce{Na2SO4}$, the sodium sulfate will ionize as you indicated:

$\ce{Na2SO4(s) ->[aq] 2Na+(aq) + SO4^{2-}(aq)}$

This reaction is essentially irreversible in that the ions in an unsaturated solution won't suddenly form solid $\ce{Na2SO4}$ again.

$\ce{HSO4-}$ is a not a weak acid, but has a $\text{pK}_\text{a2} = 1.99$ which makes a moderately strong acid. So for pH > 3 essentially all the $\ce{HSO4-}$ will ionize:

$\ce{HSO4-(aq) + H2O ->[pH > 3] H3O+ (aq) + SO4^2- (aq)}$

Thus $\ce{SO4^{2-}}$ is such a weak base that reverse reaction doesn't happen when you dissolve $\ce{Na2SO4}$, and a solution of $\ce{Na2SO4}$ in distilled water stays neutral rather than becoming basic.

$\ce{\require{cancel}H2O + SO4^2- \cancel{->} HSO4- + OH-}$

To think about a different example acetic acid is a weaker acid, $\text{pK}_\text{a} = 4.76$, so the acetate anion is a stronger base than the sulfate anion. A solution of sodium acetate would become weakly basic.

One last point - the statement should be "sodium sulfate a neutral salt in aqueous solution." In other solvents sodium sulfate could act as a base.


Here is the somewhat weak explanation given by Wikipedia:

Sodium sulfate is a neutral salt: its aqueous solutions exhibit a pH of 7. The neutrality of such solutions reflects the fact that sulfate is derived, formally, from the strong acid sulfuric acid. Furthermore, the Na+ ion, with only a single positive charge, only weakly polarizes its water ligands provided there are metal ions in solution. Sodium sulfate reacts with sulfuric acid to give the acid salt sodium bisulfate.

It definitely leaves some unanswered questions like what they mean by "metal ions in solutions": what metal ions (sodium ion?), what concentrations?

But the gist of it makes sense and makes a reasonable answer to me.


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