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It is commonly taught in chemistry class that the conjugate base of a weak acid (well, at least if it's not weaker than water) is a weak base, as shown by the relation

$$\mathrm{p}K_\mathrm{a} = 14 - \mathrm{p}K_\mathrm{b} \iff K_\mathrm{a}=\frac{K_\mathrm{w}}{K_\mathrm{b}}$$

where the $K_\mathrm{a}$ is of the conjugate acid, and $K_\mathrm{b}$ of the base. For weak monoprotic acids this means that their conjugate bases are weak, e.g. various carboxylic acids, phenol and hydrofluoric acid. For polyprotic acids, if the conjugate base's next $K_\mathrm{a}$ is smaller than its $K_\mathrm{b}$, then it would also form an alkaline solution but not fully dissociate, e.g. bicarbonate ion (carbonic acid), bisulfide ion (hydrogen sulfide), ammonia (ammonium ion).

However, according to Wikipedia as well as my teacher, sodium sulfate $\ce{Na2SO4}$ is neutral, but Wikipedia also says the acidic hydrogen on the bisulfate ion has $\mathrm{p}K_\mathrm{a} \approx 2$, so sulfate would have $\mathrm{p}K_\mathrm{a} \approx 12$ and a $1\ \mathrm{M}$ solution of sulfate with a perfectly non-acidic counterion would have $\mathrm{pH} \approx 8$. Is sodium sulfate in fact neutral? If so, what would be the reason? Could it be that $\ce{Na+}$ is slightly acidic? Or is the formula for conjugate basicity inapplicable here?

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  • $\begingroup$ Duplicate (though without answer): chemistry.stackexchange.com/questions/36913/… $\endgroup$ – a-cyclohexane-molecule Aug 16 '16 at 13:46
  • $\begingroup$ @a-cyclohexane-molecule Hm... Is it then a good idea to close this? $\endgroup$ – busukxuan Aug 16 '16 at 13:47
  • $\begingroup$ @a-cyclohexane-molecule I mean delete. Should I delete it? $\endgroup$ – busukxuan Aug 16 '16 at 13:48
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    $\begingroup$ I don't know, considering that the duplicate post has no answer. I would wait for someone more experienced to reply. $\endgroup$ – a-cyclohexane-molecule Aug 16 '16 at 13:51
  • $\begingroup$ Do not use mhchem in question title due to search- ability issues. $\endgroup$ – Nilay Ghosh Aug 20 '16 at 4:26
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Given $\text{pK}_{a_1} = -3$ and $\text{pK}_{a_2} = 1.9$ for sulphuric acid, the pH for a 1M solution is indeed expected to be $7.9$, $7.4$ for a 0.1M, and $8.0$ for a saturated solution. So it is slightly basic.

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  • $\begingroup$ So the claim of neutrality is either a very rough approximation or an oversimplification of theory? $\endgroup$ – busukxuan Aug 16 '16 at 14:12
  • $\begingroup$ It depends on your definition of neutrality. From a theoretical point of view, if you would define neutral to be the state where you have exactly equal concentrations of hydronium and hydroxide ions, you would hardly ever get a neutral solution in practice. On the other hand, it might be OK to call a solution with 6 < pH < 8 neutral for many applications. $\endgroup$ – aventurin Aug 16 '16 at 14:23
  • $\begingroup$ Wouldn't other neutral salts like $NaCl$ form pH 7.0 solutions though? $\endgroup$ – busukxuan Aug 16 '16 at 14:25
  • $\begingroup$ Since $\ce{HCl}$ is a very strong acid I whould expect that it would not affect the concentration of hydronium ions by reaction of water molecules with the chloride anions in a measurable way. However, the ionic strength of a higher concentrated $\ce{NaCl}$ solution would cause a slight deviation from pH 7. $\endgroup$ – aventurin Aug 16 '16 at 15:09
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    $\begingroup$ Yes, according to the principle of electroneutrality it must be neutral. $\endgroup$ – aventurin Aug 16 '16 at 15:38

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