Acidity of non-transition cations

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?


