thermal stability of Group 2 carbonates

Heat a Group 2 carbonate strongly and it breaks apart, releasing carbon dioxide and leaving the metal oxide behind — this is exactly how limestone is roasted into quicklime in a kiln, CaCO3 giving CaO + CO2. But how hard you have to heat it depends sharply on which metal it is: magnesium carbonate decomposes at a few hundred degrees, while barium carbonate clings to its CO2 until well over a thousand. Stability rises as you go down the group.

Why does the heavier carbonate resist decomposition? The carbonate ion CO3 2- is large, and a small, highly charged cation sitting next to it distorts and polarizes it — pulling electron density toward the metal and weakening the carbon-oxygen bonds within the carbonate, effectively prying out one oxygen as it leaves with the metal to form the oxide. Magnesium's small Mg2+ polarizes the carbonate strongly, destabilizing it, so MgCO3 decomposes easily and at low temperature. As you descend the group the cation grows, its polarizing power falls, the carbonate is left more nearly undisturbed, and so it is more thermally stable. The trend in decomposition temperature climbs steadily: MgCO3, CaCO3, SrCO3, BaCO3.

This matters because it is a clean, predictive illustration of cation polarizing power (the charge-to-size logic of Fajans' rules) applied to thermal stability, and the same reasoning explains the parallel trend in Group 2 nitrates and the easy decomposition of lithium's carbonate. The honest caveat for beginners is that the trend feels backwards: you might expect the most reactive metal at the bottom of the group to give the least stable compound, but here it is the small cation at the top, not the most reactive one at the bottom, that makes the least stable carbonate. Stability of the salt and reactivity of the metal are governed by different things.