Chemistry
The Thermodynamics of Endothermic Dissolution in Ammonium Nitrate Cold Packs
Quick fact
Ammonium nitrate cold packs can drop from room temperature to about 35°F (2°C) in under 30 seconds, absorbing roughly 26 kJ of heat per mole of salt dissolved—a dramatic demonstration that some spontaneous processes absorb heat rather than release it.
Why this is interesting
You grab a cold pack from the first-aid kit, squeeze it, and within seconds it becomes freezing. How can a chemical reaction make things colder?
Read the full explanation
Understanding The Thermodynamics of Endothermic Dissolution in Ammonium Nitrate Cold Packs
When you break the inner pouch of a cold pack, water mixes with solid ammonium nitrate (NH₄NO₃). The salt begins to dissolve, and as it does, it pulls heat from the surrounding water and the pack's exterior. The pack feels cold because thermal energy is being consumed as the salt's ionic bonds break and its ions get hydrated by water molecules. This process is called an endothermic dissolution—it absorbs heat from the environment. But here's the puzzle: many chemical changes release heat (exothermic), like burning wood. Why would a process that soaks up energy happen at all? The answer lies in the balance of energy and disorder: the salt's particles spread out into a much more disordered state, which tends to drive the process forward even though it's energetically 'uphill'.
A deeper explanation
The spontaneity of ammonium nitrate dissolution is governed by the Gibbs free energy change: ΔG = ΔH - TΔS. For the process to be spontaneous, ΔG must be negative. Here, ΔH is positive because breaking the ionic lattice requires energy (lattice energy) that is only partially compensated by the energy released when water molecules surround the ions (hydration energy). So the system absorbs heat from the surroundings. This alone would make ΔG positive and the process nonspontaneous. However, the dissolution produces a large increase in entropy (ΔS 0) because the ordered crystal structure is destroyed and the ions become freely dispersed in solution, creating far more microstates. At room temperature, the TΔS term is large enough to overcome the positive ΔH, making ΔG negative. Thus, the cold pack works because entropy wins the thermodynamic tug-of-war, allowing heat to be withdrawn from the pack and its surroundings.