Thermal stability order of NaF, MgF2 and AlF3

I came up with a question to arrange thermal stability order of $\ce{NaF}$, $\ce{MgF2}$ and $\ce{AlF3}$ and I think the answer is $\ce{NaF>MgF2>AlF3}$ because $\ce{Na+}$ has largest ionic radius among the cations(anion is same) and also NaF has the greatest ionic character. And as we know greater the ionic character, the greater is the thermal stability.


But the answer is exactly opposite that is $\ce{NaF<MgF2<AlF3}$. Where am I wrong?


Do lattice enthalpy also have a role to play in determining the thermal stability of metal chlorides and fluoride? If yes, then how?




Metal fluoride thermal stability trends are largely dictated by lattice enthalpies; what matters is the relative size of the cation in consideration (the charge/radius ratio - a measure of polarizability - is also a key factor). Smaller cations can better hold small anions like fluoride. 
The larger the lattice enthalpy factor, the tougher thermal decomposition becomes.
As atomic size decreases across a period, the order of stability is AlF3 > MgF2> NaF.


 0