Stirring salt into water involves a fundamental scientific process of energy transfer. All chemical and physical changes either release or absorb heat from their surroundings, and the dissolution of salt is no exception. The resulting temperature change can be subtle or dramatic, depending on the specific substance dissolved.
Defining Endothermic and Exothermic Reactions
The energy change that accompanies a process is categorized by how it interacts with its environment. A reaction or process is classified as endothermic if it takes in heat from the surroundings. This absorption of thermal energy causes the immediate environment, such as the water and the container, to experience a drop in temperature and feel colder.
Conversely, an exothermic process is one that releases energy, typically in the form of heat, into the surroundings. This warming effect means that the water and the container will feel warmer to the touch.
The Energy Balance of Dissolution
The thermal outcome of dissolving a salt is the result of two competing energy processes. When a solid salt crystal is placed in water, the first action requires energy input to break the strong electrostatic forces holding the positive and negative ions together in the crystal lattice structure. This necessary input of energy to separate the ions is an endothermic step known as the Lattice Energy.
The second process begins immediately after the ions separate, as the polar water molecules rush in to surround and stabilize the now-free ions. The formation of these new, attractive forces between the water molecules and the charged ions releases energy, which is an exothermic step called the Hydration Energy.
The final thermal nature of the dissolution, known as the Enthalpy of Solution, is determined by the net difference between these two energy values. If the initial energy required to break the crystal lattice (Lattice Energy) is greater than the energy released during the stabilization of the ions by water (Hydration Energy), the process has a net absorption of heat. This net absorption results in an endothermic reaction and a cooling sensation. If the energy released by the water molecules forming new attractions (Hydration Energy) is greater than the energy needed to break the salt apart, the process has a net release of heat, resulting in an exothermic reaction, causing the temperature of the water to rise.
The Thermal Nature of Common Salts
For the most common substance, sodium chloride (\(\text{NaCl}\)), or everyday table salt, the dissolution process is slightly endothermic. The energy needed to pull the sodium and chloride ions apart from the crystal lattice is marginally greater than the energy released when the water molecules surround them. This means that dissolving table salt causes a very slight drop in the water’s temperature, although it is often too small to be easily noticed without specialized equipment.
The thermal outcome is highly dependent on the specific salt’s chemical structure, meaning the answer is not the same for all ionic compounds. Salts used in instant cold packs, such as ammonium nitrate (\(\text{NH}_4\text{NO}_3\)) or potassium nitrate (\(\text{KNO}_3\)), are excellent examples of strongly endothermic dissolution. The required lattice energy for these salts far exceeds the hydration energy, causing a significant and rapid drop in temperature that makes the packs effective for treating injuries.
On the opposite end of the spectrum, salts like calcium chloride (\(\text{CaCl}_2\)) and magnesium sulfate (\(\text{MgSO}_4\)) exhibit a net exothermic dissolution. In these cases, the energy released during the hydration of the ions is substantially greater than the energy absorbed to break the crystal apart. Calcium chloride is often used in road de-icers and commercial hot packs because its highly exothermic dissolution releases enough heat to melt ice and warm the surroundings.