$\ce{CaCO3}$ is more thermally stable than $\ce{MgCO3}.$ I do understand that a larger cation would be able to stabilize a larger anion to a greater extent, but why can't we use Fajans' rule or lattice energy to explain?


1 Answer 1


(1) Thermal Stability:

$\ce{CaCO3(s) ->[\Delta] CaO(s) + CO2(g)}$

$\ce{MgCO3(s) ->[\Delta] MgO(s) + CO2(g)}$

Because of smaller $\ce{Mg^2+}$ ion, it will polarise the $\ce{CO3^2-}$ ion more than $\ce{Ca^2+}$ ion. So the removal of $\ce{CO2}$ will be easier. See this explanation.

Fajans' rule compares the covalent character in ionic compounds. For which, it compares the polarising power of cations if the anion is same. Hence, this rule in a way, explains the thermal stability of these two compounds.

(2) Fajans' Rule:

The covalent character is less in $\ce{CaCO3}$ than $\ce{MgCO3}$.

(3) Lattice Energy:

The lattice enthalpy of an ionic solid is defined as the energy required to completely separate one mole of a solid ionic compound into gaseous constituent ions.

$\ce{CaCO3(s) ->[\Delta] Ca^2+(g) + CO3^2-(g)}$

$\ce{MgCO3(s) ->[\Delta] Mg^2+(g) + CO3^2-(g)}$

$\ce{MgCO3(s)}$ has more lattice energy but it is less thermally stable than $\ce{CaCO3(s)}$. Hence, it cannot explain the trend in thermal stability because of different products obtained in (1) and (3). More on this is here.


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