My general chemistry textbook discusses how cations behave as (Brønsted-Lowry) acids in water. There's a picture of a $1.0~\mathrm M$ solution of Calcium nitrate with a pH indicator indicating a pH of around 6.9. How is it that $\ce{Ca^{2+}}$ ions behave as acids at all?

The book does explains how transition metals behave as acids: they form coordination spheres with water, and one ligand water molecule loses a proton to a free water molecule. The hydroxide is bound to the transition metal, thus not affecting pH.

However, to my knowledge, non-transition metals, such as $\ce{Ca^{2+}}$ or $\ce{Al^{3+}}$ do not form coordination compounds. Even if they can solvate some hydroxide ions, isn't the bond much weaker and therefore the hydroxide not effectively removed from solution like by a transition metal?

What would be the acid equilibrium expression for these cations?

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    $\begingroup$ Calcium actually does exactly what you're describing transition metal cations to do (binds $\ce{H2O}$ and then releases $\ce{H^+}$). Aluminum does it even better, as does magnesium. This is exactly why using $\ce{MgSO4}$ to dry solutions with highly acid-sensitive compounds is not recommended. $\endgroup$ – SendersReagent Mar 8 '16 at 3:57
  • $\begingroup$ So why is a solution of $Ca(OH)_2$ basic? It dissolves to form $OH^-$ ions, but the subsequent $H_3O^+$ released by the solvated calcium ion counteracts that to some degree. $\endgroup$ – Yunfei Ma Mar 8 '16 at 13:12
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    $\begingroup$ Calcium hydroxide forms a lot more hydroxide ions than the amount that get "lost" to binding with the calcium ions. So there are plenty of hydroxide ions left over to make limewater basic. Ditto for magnesium hydroxide solutions, albeit with much smaller solubility. $\endgroup$ – Oscar Lanzi Mar 9 '16 at 3:12

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