Why can't the Fajans' rule and lattice energy theory be used to explain the difference in thermal stability of magnesium and calcium carbonates?

$\ce{CaCO3}$ 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) 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](https://chem.libretexts.org/Bookshelves/Inorganic_Chemistry/Supplemental_Modules_and_Websites_(Inorganic_Chemistry)/Descriptive_Chemistry/Elements_Organized_by_Block/1_s-Block_Elements/Group__2_Elements%3A_The_Alkaline_Earth_Metals/1Group_2%3A_Chemical_Reactions_of_Alkali_Earth_Metals/The_Thermal_Stability_of_the_Nitrates_and_Carbonates) 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](https://chemistry.stackexchange.com/questions/68588/difference-between-lattice-energy-and-thermal-stability?rq=1).


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