Solubility of the Hydroxides, Sulfates and Carbonates of the Group 2 Elements in Water

This page looks at the solubility in water of the hydroxides, sulfates and carbonates of the Group 2 elements – beryllium, magnesium, calcium, strontium and barium. Although it describes the trends, there isn't any attempt to explain them on this page – for reasons discussed later.

You will find that there aren't any figures given for any of the solubilities. There are major discrepancies between the figures given by two common UK A-level Data Books (Nuffield Advanced Science Book of Data, and Chemistry Data Book by Stark and Wallace). There are also important inconsistencies within the books (one set of figures doesn't agree with those which can be calculated from another set). I haven't been able to find data which I am sure is correct, and therefore prefer not to give any.

This is a trend which holds for the whole Group, and applies whichever set of data you choose.

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". 1 litre of pure water will dissolve about 1 gram of calcium hydroxide at room temperature.

Barium hydroxide is soluble enough to be able to produce a solution with a concentration of around 0.1 mol dm^-3 at room temperature.

The simple trend is true provided you include hydrated beryllium sulfate in it, but not if the beryllium sulfate is anhydrous.

The Nuffield Data Book quotes anyhydrous beryllium sulfate, BeSO₄, as insoluble (I haven't been able to confirm this from any other source), whereas the hydrated form, BeSO₄.4H₂O is soluble. (The Data Books agree on this – giving a figure of about 39 g dissolving in 100 g of water at room temperature.)

Figures for magnesium sulfate and calcium sulfate 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 sulfuric acid to give lots of hydrogen and a colourless solution of magnesium sulfate. Notice that you get a solution, not a precipitate. The magnesium sulfate is obviously soluble.

You may also remember that barium sulfate is formed as a white precipitate during the test for sulfate ions in solution. The ready formation of a precipitate shows that the barium sulfate must be pretty insoluble. In fact, 1 litre of water will only dissolve about 2 mg of barium sulfate 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 0.02 g per 100 g of water at room temperature.

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.

There are no simple examples which might help you to remember the carbonate trend.

Before I started to write this page, I thought I understood the trends in solubility patterns including the explanations for them. The more I have dug around to try to find reliable data, and the more time I have spent thinking about it, the less I'm sure that it is possible to come up with any simple explanation of the solubility patterns.