Ammonium nitrate is one of the classic examples of an endothermic dissolution. People will usually explain why it's soluble by stating that the spontaneous dissolution process of ammonium nitrate is what drives the reaction, meaning that the obvious answer is it's soluble because ΔG of dissolving it is negative. Here is an example.
However, after thinking a bit more about it, this explanation seems very vague and doesn't explain it well. Enthalpy usually has a stronger role in determining the solubility of salts, since their ΔH values are usually much larger than the TΔS values they would have, especially at room temperature.
Both calcium and aluminum fluoride are also salts that have an endothermic dissolution, but by the number of ions dissociated, one would think that ΔS of dissolving both of these salts would be much higher than that of ammonium nitrate. Ammonium nitrate is still several times more soluble though.
My question is, what mechanisms allow ammonium nitrate to have such a high ΔS when dissolving in solution that allow it to still be very soluble, despite having an endothermic dissolution?