Why does CaSO4 break up when adding HCl but not BaSO4?

Question

I observed in a solution with CaSO4 that the white precipitate dissolved when I added concentrated hydro chloric acid (37%). I assume CaSO4 had a divorce and Ca decided to hook up with two Cl- ions, forming CaCl2.

Having access to the big sister of Ca, Barium, I tried to stress test the marriage of BaSO4 in a similar way by enticing Ba to hook up with Cl- ions. But it seems that Barium is a very committed partner, and the precipitate could not be dissolved.

My question is: why is this difference observed, and what do I need to do to obtain the same result with Ba?

I saw a similar question on this forum with Ca (CO3), HCl and H2SO4 that were answered by mentioning an equilibrium and shifting the reaction to one side. So I assume the general answer to my question could be that it is easier to reverse the reaction with Ca than with Ba. If that is the case, would vigorous shaking + high temperature + high concentration of H+ ions help to convince Ba to consider Cl- ions instead? Or is BaSO4 one of those "game-over" reactions, there is no going back once you get them.

Thank you very much for your time

Answer 1

Sulphates of alkali earth metals dissolve like this:

$$\ce{MSO4(s) <=>[H2O] M^2+(aq) + SO4^2-(aq)}\tag{1}$$

with the solubility product constant (kept original eq. numbering):

$$K_\mathrm{sp} = a(\ce{M^2+})a(\ce{SO4^2-}),\tag{3}$$

where $a(\ce{X})$ is thermodynamic activity of the ion $\ce{X}$. For diluted solutions is can be approximated by molar concentrations (denoted as []) as

$$K_\mathrm{sp} = [\ce{M^2+}][\ce{SO4^2-}]\tag{4}.$$

• $K_\mathrm{sp} \gt [\ce{M^2+}][\ce{SO4^2-}] \implies \text{net dissolution}$
• $K_\mathrm{sp} \lt [\ce{M^2+}][\ce{SO4^2-}] \implies \text{net precipitation}$

As a hydrogensulphate anion is not strong acid ($\mathrm{p}K_\mathrm{a}=1.99$) , concentrated hydrochloric acid protonizes in large extent sulphate to hydrogensulphate.

$$\ce{SO4^2-(aq) + H+(aq)<=>>[conc. HCl] HSO4-(aq)}\tag{2}$$

Decreasing of sulphate concentration disbalances dissolution equilibrium and there is ongoing net dissolution until the equilibrium is reached again.

$\ce{CaSO4}$ is slightly soluble ($\pu{0.26 g/100 ml}$ at $\pu{25 ^{\circ}C}$ (dihydrate)), therefore well soluble in $\ce{HCl}$.

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Solubility of $\ce{BaSO4}$ is much less, ($\pu{0.2448 mg/100 mL}$ at $\pu{20 ^{\circ}C}$ ) so it needs (hot) concentrated $\ce{H2SO4}$. Concentrated $\ce{H2SO4}$ acts here as the solvent. $\pu{2-4\%}$ of $\ce{H2O}$ is practically all converted to $\ce{H3O+(solv)}$. Additionally, $\ce{H2SO4}$ partially autodissociates.

\begin{align} \ce{H2SO4(l) + H2O &-> HSO4-(solv) + H3O+(solv)}\tag{5}\\ \ce{2 H2SO4(l) &<=> HSO4-(solv) + H3SO4+(solv)}\tag{6} \end{align}

This causes extremely high activity of solvated $\ce{H+}$, extremely low activity of $\ce{SO4^2-(solv)}$,

\begin{align} \ce{SO4^2-(solv) + H3O+(solv) &-> HSO4-(solv) + H2O(solv)}\tag{7}\\ \ce{SO4^2-(solv) + H2SO4(l) &-> 2 HSO4-(solv)}\tag{8}\\ \ce{SO4^2-(solv) + H3SO4+(solv) &-> HSO4-(solv) + H2SO4(l)}\tag{9} \end{align}

leading to relatively high $\ce{BaSO4(s)}$ solubility.

It can be also described as

$$\ce{BaSO4(s) + H2SO4(l) <=>[H2SO4] Ba(HSO4)2(solv)},$$

in a way analogical to:

$$\ce{CaCO3(s) + H2O(l) + CO2(aq)) <=> Ca(HCO3)2(aq)}$$