# Neutral anions in salts and strong acids are similar - is it a coincidence? [closed]

Recently I was studying salts and my professor provided a list of anions which are considered neutral in salts (that is, they do not change the $\mathrm{pH}$ of an aqueous solution when added as a salt): $\ce{Cl-}$, $\ce{Br-}$, $\ce{I-}$, $\ce{NO_3^-}$, $\ce{ClO_4^-}$.

Now I'm looking at this and I realized the list seems to be almost identical to a list of strong acids: $\ce{HCl}$, $\ce{HBr}$, $\ce{HI}$, $\ce{HNO_3}$, $\ce{HClO_4}$, $\ce{H_2SO_4}$.

Is this a coincidence? If it's not, what's the reason, and why is $\ce{SO_4^2-}$ not on the list of neutral ions?

## closed as off-topic by airhuff, Mithoron, a-cyclohexane-molecule, M.A.R. ಠ_ಠ, aventurinMar 18 '18 at 15:52

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• The conjugate acid of a weak base is a strong acid, and the anions you have mentioned are weak bases. – a-cyclohexane-molecule Mar 18 '18 at 6:21

Though couched in different terms, the first part this question is essentially the same as this other one:

Does a strong acid's conjugate base act as Brønsted-Lowry base?

No, this is not a coincidence. The conjugate base of a strong acid will always be a weak base. The 'neutral salt anions' of your professor's nomenclature are just weak inorganic Brönsted bases.

The reason why $$\ce{SO4^2-}$$ doesn't appear in this list of 'neutral salt anions' is because it actually can noticeably affect the $$\mathbf{pH}$$ of a solution when added as a salt. $$\ce{H2SO4}$$ is often treated as though it's a diprotic strong acid, but in reality only the first deprotonation is strongly acidic. This can be seen by examining its acid dissociation constants; from Wikipedia:

$$\mathrm pK_{\mathrm a1} = -3$$
$$\mathrm pK_{\mathrm a2} = 1.99$$

This $$\mathrm pK_{\mathrm a2}$$ is about the same as the first $$\mathrm pK_{\mathrm a1}$$ of $$\ce{H3PO4}$$ $$($$$$2.148$$$$)$$, and is actually somewhat less acidic than nitroacetic acid $$($$$$1.68$$$$)$$.

So, in some cases sulfate may act more or less as a 'neutral salt anion', when the $$\mathrm{pH}$$ is sufficiently high that it will all stay doubly deprotonated, but in others it can definitely influence the solution $$\mathrm{pH}$$. I've actually used sulfate/bisulfate before as an acidic buffer for various projects at work—works great!