hydrolysis of aqua ions

Dissolve aluminium chloride or iron(III) nitrate in pure water and measure the pH — it is distinctly acidic, even though no obvious acid was added. The cause is one of the quietest acids in chemistry: the metal cation itself, surrounded by its tightly bound water molecules. A small, highly charged metal ion does not float around bare; it is an aqua ion, for example [Al(H2O)6]3+ or [Fe(H2O)6]3+, an octahedral cluster of six water molecules coordinated to the metal. Those coordinated waters are surprisingly acidic.

Here is the mechanism. The metal ion's positive charge pulls electron density away from the oxygen of each bound water, which in turn weakens the O-H bonds of that water and makes its protons easier to lose. So the aqua ion behaves as a Bronsted acid, donating a proton from a coordinated water to a free water molecule: [Al(H2O)6]3+ + H2O gives [Al(H2O)5(OH)]2+ + H3O+. This is simultaneously a Lewis-acid story — the hard, high-charge Al3+ polarizes the O-H bond — and a Bronsted one. The effect grows with charge and shrinks with size: the higher the charge density (charge divided by radius), the more acidic the aqua ion. So 3+ ions like Al3+, Fe3+, and Cr3+ are noticeably acidic (pKa of [Fe(H2O)6]3+ is around 2, comparable to a weak organic acid), 2+ ions are mildly acidic, and big 1+ ions like Na+ and K+ are essentially non-acidic.

Aqua-ion hydrolysis explains a great deal of real solution chemistry. It is why solutions of Al3+ and Fe3+ salts are acidic and why they precipitate gelatinous hydroxides as the pH is raised; it underlies the formation of polynuclear hydroxo-bridged species and ultimately metal-oxide gels and rust; it governs which metal ions stay dissolved in natural waters and at what pH; and it is the first step toward understanding why high-oxidation-state metals exist in water as oxo-anions (like CrO4 2- or MnO4-) rather than as simple cations — successive loss of protons from aqua ligands, taken to the limit, replaces water with oxide.