The oxides of calcium, strontium, and barium are basic and the hydroxides are strongly basic. Group II metal oxides become more basic as you go down the column. This trend is easily seen if you compare the electronegativity of the group II metal to the electronegativity of oxygen.
As you can see the electronegativities of the metals decrease down the column making the change in electronegativities increases down the group. The greater the difference in electronegativity the more ionic the metal-oxygen bond becomes. The more ionic the metal-oxygen bond the more basic the oxide is. Group II metal hydroxides become more soluble in water as you go down the column. This trend can be explained by the decrease in the lattice energy of the hydroxide salt and by the increase in the coordination number of the metal ion as you go down the column.
The larger the lattice energy the more energy it takes to break the lattice apart into metal and hydroxide ions. Since the atomic radii increase down the group it makes sense that the coordination numbers also increases because the larger the metal ion the more room there is for water molecules to coordinate to it. This simple trend is true provided hydrated beryllium sulfate is considered, but not anhydrous beryllium sulfate.
Solubility figures for magnesium sulfate and calcium sulfate also vary depending on whether the salt is hydrated or not, but the variations are less dramatic. The carbonates become less soluble down the group.
All the Group 2 carbonates are very sparingly soluble. Magnesium carbonate, for example, has a solubility of about 0. There is little data for beryllium carbonate, but as it reacts with water, the trend is obscured. The trend to lower solubility is, however, broken at the bottom of the group: barium carbonate is slightly more soluble than strontium sulfate. There are no simple examples of this trend. Magnesium hydroxide appears to be insoluble in water. However, if you shake it with water, filter it and test the pH of the solution, you find that it is slightly alkaline.
This shows that there are more hydroxide ions in the solution than there were in the original water. Some magnesium hydroxide must have dissolved. Calcium hydroxide solution is used as "lime water". Barium hydroxide is soluble enough to be able to produce a solution with a concentration of around 0. The simple trend is true provided you include hydrated beryllium sulphate in it, but not if the beryllium sulphate is anhydrous.
The Data Books agree on this - giving a figure of about 39 g dissolving in g of water at room temperature. Figures for magnesium sulphate and calcium sulphate also vary depending on whether the salt is hydrated or not, but nothing like so dramatically. You are probably familiar with the reaction between magnesium and dilute sulphuric acid to give lots of hydrogen and a colourless solution of magnesium sulphate. Notice that you get a solution, not a precipitate.
The magnesium sulphate is obviously soluble. You may also remember that barium sulphate is formed as a white precipitate during the test for sulphate ions in solution. The ready formation of a precipitate shows that the barium sulphate must be pretty insoluble. In fact, 1 litre of water will only dissolve about 2 mg of barium sulphate at room temperature. None of the carbonates is anything more than very sparingly soluble.
Magnesium carbonate the most soluble one I have data for is soluble to the extent of about 0. I can't find any data for beryllium carbonate, but it tends to react with water and so that might confuse the trend.
The trend to lower solubility is, however, broken at the bottom of the Group. Barium carbonate is slightly more soluble than strontium carbonate. Before I started to write this page, I thought I understood the trends in solubility patterns including the explanations for them.
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