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








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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?




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?](https://chemistry.stackexchange.com/questions/67685/does-a-strong-acids-conjugate-base-act-as-br%C3%B8nsted-lowry-base)


**No, this is not a coincidence.** The [conjugate base](https://en.wikipedia.org/wiki/Conjugate_acid) 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.




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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](https://en.wikipedia.org/wiki/Sulfuric_acid):



> 
> $\mathrm pK\_{\mathrm a1} = -3$  
> 
> $\mathrm pK\_{\mathrm a2} = 1.99$
> 
> 
> 


This $\mathrm pK\_{\mathrm a2}$ is about the same as the first $\mathrm pK\_{\mathrm a1}$ of $\ce{H3PO4}$ $($[$2.148$](https://en.wikipedia.org/wiki/Phosphoric_acid)$)$, and is actually somewhat *less* acidic than nitroacetic acid $($[$1.68$](https://en.wikipedia.org/wiki/Nitroacetic_acid)$)$.


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!


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