Why is Barium Phosphate soluble in dil HCl? [closed]

Why is Barium Phosphate soluble in dil HCl While Barium Sulphate is not?

• What do you know about the acid-base equilibria? More to the point, what dissolved anion is in equilibrium with solid barium phosphate and what happens to it as you add an acid? – Ivan Neretin May 17 at 18:46
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Many phosphates are not soluble in water at standard temperature and pressure, except for the sodium, potassium, rubidium, cesium, and ammonium phosphates, which are all water-soluble. As a rule, the hydrogen and dihydrogen phosphates are slightly more soluble than the corresponding phosphates (Wikipedia). For example, based on Wikipedia Solubility Table, the solubility of $$\ce{BaHPO4}$$ is $$\pu{0.013 g}$$ in $$\pu{100 mL}$$ at $$\pu{20 ^{\circ}C}$$, and a CDC.gov Report states that:

The water solubility of barium compounds increases with decreasing $$\mathrm{pH}$$.

Nevertheless, the solubility studies of the phosphates of strontium, barium, and magnesium have been done previously (Ref.1). Alkaline earth phosphates as a whole have been reviewed as well (Ref.2).

That said, let's look at phosphates behaviour with $$\mathrm{pH}$$. The three $$\mathrm{p}K_\mathrm{a}$$ values for phosphoric acid are 2.16, 7.21, and 12.32. Thus, aqueous phosphate exists in four forms (Wikipedia):

• In strongly basic conditions, the phosphate ion ($$\ce{PO4^3-}$$) predominates.
• In weakly basic conditions, the hydrogen phosphate ion ($$\ce{HPO4^2-}$$) is prevalent.
• In weakly acidic conditions, the dihydrogen phosphate ion ($$\ce{H2PO4-}$$) is most common.
• In strongly acidic conditions, trihydrogen phosphate ($$\ce{H3PO4}$$) is the main form.

When dilute $$\ce{HCl}$$ (say $$\pu{3 M}$$ and hence $$\mathrm{pH} \approx -0.48$$) is added to solid $$\ce{Ba3(PO4)2}$$, following reactions will occur with aqueous hydronium ions: $$\ce{Ba3(PO4)2(s) + nH2O(l) <=> 3Ba^2+(aq) + 2PO4^3-(aq)}$$ $$\ce{PO4^3-(aq) + H3O+(aq) -> HPO4^2-(aq) + H2O(l)}$$ $$\ce{HPO4^2-(aq) + H3O+(aq) -> H2PO4-(aq) + H2O(l)}$$ $$\ce{H2PO4-(aq) + H3O+(aq) -> H3PO4(aq) + H2O(l)}$$

Because, $$\mathrm{pH}$$ of the acid solution is lower than the $$\mathrm{p}K_\mathrm{a1}$$ of the $$\ce{H3PO4}$$ acid, $$\ce{Ba3(PO4)2}$$ solid would slowly dissolve to give $$\ce{H3PO4(aq)}$$ and $$\ce{BaCl2(aq)}$$ in the solution (Channel Le Chatelier's principle).

However, this phenomenon would not happen with $$\ce{BaSO4}$$ solid because $$\mathrm{p}K_\mathrm{a1}$$ of $$\ce{H2SO4}$$ is about $$-3$$, which is much lower than the operating $$\mathrm{pH}$$. Therefore, since $$\mathrm{p}K_\mathrm{a2}=1.99$$ for $$\ce{H2SO4}$$, it is possible to form $$\ce{Ba(HSO4)2}$$, yet its solubility is said to be low as well ($$\approx \pu{0.015 g}$$ in $$\pu{100 mL}$$ of water at $$\pu{25 ^{\circ}C}$$).

References:

1. L. E. Holt, Jr., J. A Pierce, C. N Kajdi, “The solubility of the phosphates of strontium, barium, and magnesium and their relation to the problem of calcification,” Journal of Colloid Science 1954, 9(5), 409–426 (https://doi.org/10.1016/0095-8522(54)90029-X).
2. R. W. Mooney, M. A. Aia, “Alkaline Earth Phosphates,” Chem. Rev. 1961, 61(5), 433–462 (DOI: 10.1021/cr60213a001).
• Minor point. Is there actually $\ce{(NH4)3PO4}$, or must ammonium phosphate be an acid salt? The normal phosphate ion "should be" strongly basic enough to displace $\ce{NH3}$. – Oscar Lanzi May 17 at 22:56
• @Oscar Lanzi: Honestly, I did not know what is the $\mathrm{p}K_\mathrm{a}$ of $\ce{(NH4)3PO4}$. But I found this infor for $\ce{(NH4)H2PO4}$: A solution of stoichometric monoammonium phosphate is acidic (pH 4.7 at 0.1% concentration, 4.2 at 5%). – Mathew Mahindaratne May 17 at 23:38