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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, 2016 at 13:46
  • $\begingroup$ @a-cyclohexane-molecule Hm... Is it then a good idea to close this? $\endgroup$
    – busukxuan
    Aug 16, 2016 at 13:47
  • $\begingroup$ @a-cyclohexane-molecule I mean delete. Should I delete it? $\endgroup$
    – busukxuan
    Aug 16, 2016 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, 2016 at 13:51
  • $\begingroup$ Do not use mhchem in question title due to search- ability issues. $\endgroup$
    – Nilay Ghosh
    Aug 20, 2016 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, 2016 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, 2016 at 14:23
  • $\begingroup$ Wouldn't other neutral salts like $NaCl$ form pH 7.0 solutions though? $\endgroup$
    – busukxuan
    Aug 16, 2016 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, 2016 at 15:09
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    $\begingroup$ Yes, according to the principle of electroneutrality it must be neutral. $\endgroup$
    – aventurin
    Aug 16, 2016 at 15:38

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