[LHM]

The limiting enthalpy of solution of ammonium nitrate is +6.6 kJ/mol at 25°C, so dissolving ammonium nitrate is endothermic. So the enthalpy of solution is the lattice enthalpy plus the enthalpy of hydration when the gaseous ions become part of the solution again. However, ammonium nitrate has big, singly charged anions and hence low lattice enthalpies, while its hydration enthalpy is probably pretty exothermic because water can form hydrogen bonds with the nitrate ions. So how come the enthalpy of solution is positive if the enthalpy of hydration is probably ore negative than the lattice enthalpy?

[rabolisk]

The enthalpy of hydration is not greater than the lattice enthalpy, which is why the enthalpy of solution of ammonium nitrate is positive. Your theories are merely theories; apparently the data contradicts it.

[Grundalizer]

Funny, I asked my advanced inorganic professor this yesterday, and couldn't find the enthalpy of solvation of ammonium nitrate. 

We studied NaCl as the example, and that had an enthalpy of solvation of 4kJ/mol, yet water doesn't get "cold" when you dissolve salt in it, I am surprised to hear ammonium nitrate's is so low???  6.6kj/mol?  That's hardly that much more than NaCl, yet it's used in cold packs?

Is there something else going on here?