In fact, the heats of solution are the experimentally measured quantities, and it is only with their help that we can get reasonable estimates of heats of hydration. The lattice energies are on somewhat firmer theoretical grounds. These "small" heats of solution are large enough to have obvious physical consequences. The cooling when ordinary table salt is dissolved in water is small, but can be felt if one makes a concentrated solution and uses an aluminum tumbler. Ammonium chloride absorbs so much heat when it dissolves that hoar frost may form on the outside of the beaker. In contrast, sodium hydroxide generates so much heat that the mixing beaker may become too hot to touch.

To a certain extent, we can account for the trends in heats that we see in the table. Ammonium chloride has a weaker lattice energy than NaCl, because the NH⁴⁺ ion is larger than Na⁺ and the binding attractions in the crystal are weaker. Unfortunately, the hydration energy also decreases with increasing ionic size, and it is difficult to predict whether lattice energy or hydration energy will show the greater change with larger ions.

The heats of hydration of Cl^- and OH^- ions are similar, so in the comparison of NaOH with NaCl, the dominant effect comes from weaker crystal forces of NaOH in comparison with NaCl. The crystal structure of NaOH is in fact a badly distorted NaCl structure, with the distortion probably arising from the fact that the OH^- ions are nonspherical. It is possible that this distortion makes the NaOH lattice easier to pull apart.