The lattice energy \(^{*}\) of Nal is \(-686 \mathrm{kJ} / \mathrm{mol}\), and the enthalpy of hydration is \(-694 \mathrm{kJ} / \mathrm{mol} .\) Calculate the enthalpy of solution per mole of solid NaI. Describe the process to which this enthalpy change applies.

Lattice energy is a measure of the strength of the forces between the ions in an ionic solid. Imagine it as the ‘glue’ holding the lattice together. For a compound like sodium iodide (NaI), the lattice structure is comprised of alternating positive (Na+) and negative (I-) ions.
To dissolve the compound in water, this strong ionic lattice must first be broken apart, which requires energy. In the NaI example, the lattice energy is given as -686 kJ/mol. Since this value is negative, it indicates that energy is released when the lattice forms. Conversely, breaking the lattice apart would require the same amount of energy but in a positive form.
Understanding lattice energy is crucial because it helps predict the solubility of a compound, its melting point, and its overall stability. The higher the lattice energy, the more stable the compound and the more energy is required to break it down.

The enthalpy of hydration refers to the energy change when gaseous ions dissolve in water and become surrounded by water molecules. It is a good indicator of how each ion interacts with water.
In the case of NaI dissolving, both sodium (Na+) and iodide (I-) ions are being hydrated. The given value, -694 kJ/mol, means that more energy is released when these ions interact with water than was used to break the ionic lattice. This is because forming new ion-water attractions can be highly exothermic, as water molecules stabilize the ions.
As an additional perspective, the strength of ion-water interactions can vary depending on the size and charge of the ions. Smaller and more highly charged ions tend to have larger enthalpies of hydration. This balancing act between breaking the lattice and hydrating the ions is what governs the solubility of substances.

Thermochemistry is the branch of chemistry that deals with the energy changes during chemical reactions. In every chemical process, energy is either absorbed or released, which correlates with changes in enthalpy—one of the main components of thermochemistry.
When discussing the dissolution of an ionic compound, thermochemistry revolves around the enthalpy changes during the breaking of the ionic lattice and the subsequent hydration of the ions. By analyzing these enthalpy changes, we can determine if the process will be endothermic (absorbing heat) or exothermic (releasing heat).
Thermochemical calculations often involve Hess’s Law, which states that the total enthalpy change for a reaction is the same, no matter how it is achieved. This principle allows us to use provided enthalpy values, such as lattice energy and enthalpy of hydration, to determine the overall enthalpy of solution.

An exothermic process is one that releases energy to the surrounding environment, usually in the form of heat.
Dissolving NaI in water, as seen in the given exercise, is an exothermic process because the overall enthalpy change (-1380 kJ/mol) is negative. This tells us that more energy is released by forming new bonds during hydration than is used to break the ionic lattice. Thus, the solution and the surroundings get warmer during this dissolution.
Recognizing an exothermic process is important for various applications. In industrial settings, exothermic reactions can be harnessed to produce heat for other purposes, while, in a biological context, they play a role in cellular energy production. Awareness of energy changes is also necessary for making predictions about reaction spontaneity and safety implications.