Dissolution involves a change in energy classified as either endothermic or exothermic. An endothermic process absorbs heat from its surroundings, causing the mixture’s temperature to decrease. Conversely, an exothermic process releases heat, leading to a noticeable temperature increase. The outcome depends entirely on the specific chemical interactions between the solute (the substance being dissolved) and the solvent (the dissolving medium). The overall heat change, measured as the enthalpy of solution (\(\Delta H_{soln}\)), determines the thermal outcome.
The Three Energy Steps of Dissolution
The overall energy change during dissolution is the net result of three distinct energy steps. These steps are viewed as a theoretical sequence, even though they happen simultaneously at the molecular level. The first two steps require an input of energy (endothermic) and involve breaking the attractive forces that hold the original components together.
The first step requires energy to separate the solute particles from each other, overcoming forces like the lattice energy in a crystal or the intermolecular forces between molecules. The second step is also endothermic, demanding energy to separate the solvent particles to create space for the solute. For a liquid solvent like water, this energy is needed to disrupt the existing intermolecular attractions, such as hydrogen bonds.
The final step involves the separated solute and solvent particles coming together to form the solution, a process called solvation (or hydration when water is the solvent). This step is always exothermic because energy is released as new, stable attractive forces form between the solute and solvent molecules. The overall enthalpy of solution is the algebraic sum of the energy consumed in the first two steps and the energy released in the third step.
When Dissolution Absorbs Heat (Endothermic Process)
Dissolution is endothermic when the energy absorbed in the first two steps is greater than the energy released during solvation. The net energy change is positive, and the system draws heat from the environment to make up the deficit. The result is that the solution and its container feel cool to the touch.
This cooling effect is harnessed in practical applications like instant cold packs, which often contain salts such as ammonium nitrate (\(\text{NH}_4\text{NO}_3\)) or potassium chloride (\(\text{KCl}\)). These substances are endothermic because they possess a high lattice energy, requiring a large amount of energy to pull the ionic crystal structure apart. The energy released when the ions are solvated by water is not sufficient to compensate for this initial energy investment.
When Dissolution Releases Heat (Exothermic Process)
Dissolution is exothermic when the energy released during the formation of new solute-solvent interactions outweighs the energy required to separate the original particles. This results in a negative overall enthalpy of solution, and the excess energy is transferred to the surroundings as heat. Consequently, the solution’s temperature rises.
A common example of this is dissolving solid sodium hydroxide (\(\text{NaOH}\)) pellets in water or the dilution of concentrated sulfuric acid (\(\text{H}_2\text{SO}_4\)). In these cases, the solvation of the ions or molecules is so energetically favorable that it easily overcomes the energy needed to break the initial bonds. The significant energy release from the powerful hydration of small, highly charged ions is the driving factor, making the process intensely exothermic.