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Why are magnesium ascorbate and potassium acetate, which are salts of strong bases with weak acids, acidic?

I discovered with litmus paper that magnesium ascorbate is acidic and found out from https://foodb.ca/compounds/FDB015417 that potassium acetate should also be acidic.

How can this be? Shouldn’t conjugate bases of weak acids, ascorbate and acetate in this case, be basic and the strong dissociated spectator ions, $\ce{K+}$ and $\ce{Mg^2+}$, neutral?

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    $\begingroup$ The site you mentioned has given $\mathrm{p}K_\mathrm{a} = 4.54$ for $\ce{CH3CO2K}$. That is actually $\mathrm{p}K_\mathrm{a}$ of acetic acid. As we all know, $\ce{CH3CO2K}$ is a slightly basic salt. $\endgroup$ Mar 24, 2021 at 0:19
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    $\begingroup$ If your potassium acetate were acidic by actual measurement, I would guess that it was not pure. If your source for magnesium ascorbate is a food supplement, I would question its purity. Perhaps you could try a recrystallization. $\endgroup$ Mar 24, 2021 at 14:27

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PubChem Magnesium ascorbate $(\ce{Mg(C6H7O6)2})$ shows a monobasic salt. Wikipedia says pKa's are 4.10 (first), 11.6 (second)

In addition to Poutnik's answer, the crystal ionic radii in pm according to wikipedia of

$$\ce{Mg^2+ is 86pm and Ni^2+ is 83pm}$$

This following coordination complex formation reaction is known $$\ce{Ni^2+(aq) + 2(dmgH2)->Ni(dmgH)2(s) + 2 H+(aq)}$$

Nickel(II) is square planar.[2] It is surrounded by two equivalents of the conjugate base (dmgH−) of dimethylglyoxime (dmgH2). (https://en.wikipedia.org/wiki/Nickel_bis(dimethylglyoximate))

Four five member rings center at the Nickel(II) ions.$$\phantom{1}$$ Similarly, after this reaction $$\ce{Mg(HA)2(s) ->[H2O] [MgA2]^2-(aq) + 2 H+(aq)}$$

Rings could center at the magnesium $$\ce{[(C4H6O4)-O^--C=CO]Mg[OC=CO^--(C4H6O4)](aq)}$$

Therefore, this complex formation reaction could be possible $$\ce{Mg(HA)2(s) ->[H2O] [MgA2]^2-(aq) + 2 H+(aq)}$$ due to

  1. similar ionic radii
  2. existing coordination complex
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$\mathrm{p}K_\mathrm{a}$ relation in the answer of Kav

Wikipedia says pKa's are 4.10 (first), 11.6 (second)

gives the hint a monobasic salt with alkali metals should be very mildly alkalic, as pKa2 is farther from pH 7.

$pH \approx \frac{\mathrm{p}K_\mathrm{a1} + \mathrm{p}K_\mathrm{a2}}{2} = \frac{4.10+11.6}{2}=7.85$

OTOH, it will probably form a complex with Mg, affecting acidity of the 2nd hydrogen. Like

$$\ce{2 H2A(aq) + Mg^2+(aq) -> [MgA2]^2-(aq) + 4 H+(aq)}$$

respectively

$$\ce{Mg(HA)2(s) ->[H2O] [MgA2]^2-(aq) + 2 H+(aq)}$$

See also citrate or EDTA complexes for this effect. AFAIK, before suitable metalochrome indicators came, metal content with excess of Na2H2-EDTA had been sometimes determined by back titration using NaOH.

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