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We use carbonates in a wide range of industrial processes including: pharmaceuticals, brewing, paper and glass manufacture and water softening. Carbonates are also a major component of many types of sedimentary rock including limestone and marble. They are also the main ingredient in the sand of beaches and in the shells of molluscs. They are also an important part of our diet.
The carbonate ion has three oxygen atoms attached to a central carbon in its Lewis structure and has the symmetrical trigonal planar geometry of carbonate. It is the conjugate base of carbonic acid, H2CO3. The CO2 ion can be removed from the chemical by heating the compound. This releases the oxygen and produces carbon dioxide gas. The molecule is then neutral again and can be reused.
As you go down Group 1, the carbonates become less soluble. This is due to the fact that the carbonate ions are less polarised by singly charged positive ions, and you need more heat to persuade them to break free of the metal oxide they’re trapped in.
Most of the metal carbonates decompose on heating to give the oxide and carbon dioxide gas, with the exception of lithium carbonate which remains solid at Bunsen temperatures. The other Group 1 carbonates do decompose at higher temperatures, but not at Bunsen temperatures.