One can explain the behavior of the oxygen compounds of second shell atoms by imagining what would happen if the positive ions, , , through , were dropped into water.

This is only an imaginary experiment, because , , , and ions, with their entire second electron shell stripped off, are too reactive to exist in solution. Nevertheless, the products that these ions would form with water are what actually are observed in solution when the oxides of these elements are dissolved in water.

A ion in aqueous solution surrounds itself with four water molecules and exists peaceably as a ion, as shown.

In an acid solution (excess of protons), beryllium also occurs as a hydrated ion. Beryllium is more electronegative than lithium, and the beryllium ion is more highly charged. It therefore pulls on the lone electron pairs of water oxygens more than does.

The bond between and a water oxygen is mainly electrostatic, but the bond between and water is partially covalent. As pulls on the water lone pairs, it weakens the O-H bonds of the water molecules.

In acidic solution merely holds on tightly to its four hydrating molecules. In basic solutions, where protons are scarce, the weakened water molecules around each can release one proton, so the hydrated ion in basic solution is instead of .

Each ion then is surrounded, not by four neutral water molecules, but by four negative hydroxide ions. The cluster, or complex ion, has a negative charge. At intermediate acidities, less than four of the molecules can lose protons.