Sodium acetate, a common salt derived from acetic acid, is frequently used in food preservation and chemistry demonstrations. Its interaction with water involves a transfer of heat, which can either warm or cool the surrounding environment. Understanding this requires defining two fundamental concepts: A process is endothermic if it absorbs heat from its surroundings, often causing a temperature drop. Conversely, a process is exothermic if it releases heat, causing the temperature to rise.
The Core Concept: Heat of Dissolution
When the common form, sodium acetate trihydrate (\(\text{NaC}_2\text{H}_3\text{O}_2 \cdot 3\text{H}_2\text{O}\)), is dissolved in water, the process is endothermic. The dissolving salt actively pulls thermal energy from the water and the container. As a result of this absorption, the solution’s temperature drops measurably, making the outside of the container feel cool to the touch. This form is most often encountered when discussing its use in phase change experiments.
Understanding the Energy Balance
The thermal outcome of dissolution is governed by two competing energy processes. The first process is endothermic: the energy required to break apart the ionic bonds holding the solid crystal lattice structure together, freeing the individual sodium and acetate ions. For sodium acetate trihydrate, the heat of solution is positive, measuring approximately \(+19.7 \text{ kJ/mol}\), confirming net heat absorption.
The second process is exothermic: the hydration energy released when the free ions form attractive interactions with surrounding water molecules. These ion-dipole forces stabilize the ions in the solution, releasing energy. The overall heat of dissolution is the sum of these two opposing energy changes. For sodium acetate trihydrate, the energy input needed to break the lattice is greater than the energy released during hydration, resulting in a net endothermic result.
A contrasting thermal behavior is observed if anhydrous sodium acetate (\(\text{NaC}_2\text{H}_3\text{O}_2\)), which lacks water molecules in its crystal structure, is dissolved. Dissolving the anhydrous form is an exothermic process, releasing heat into the solution. This difference occurs because the trihydrate crystal already incorporated water, meaning a portion of the total hydration energy was released during its formation. The anhydrous form thus has a higher potential for hydration energy release, which overcomes its lattice energy, resulting in a negative heat of solution of about \(-17.1 \text{ kJ/mol}\).
The Reverse Reaction: Crystallization
While the dissolution of sodium acetate trihydrate is endothermic, the substance is famous for its exothermic behavior during the reverse process: crystallization. Crystallization occurs when the dissolved ions bond back together to form a solid crystal structure. Since this is the opposite of endothermic dissolution, the crystallization process must release an equivalent amount of thermal energy.
This energy release is typically demonstrated using a supersaturated solution of sodium acetate, commonly known as “hot ice.” This solution contains more dissolved solute than is normally stable at that temperature. The solution can be cooled significantly below the crystallization temperature, a state called supercooling, without forming a solid. When a small physical disturbance, such as adding a seed crystal or clicking a metal trigger, is introduced, it provides a nucleation site for the ions to rapidly reorganize. This instant formation of the solid lattice structure releases the stored latent heat, causing the temperature to quickly rise to the trihydrate’s melting point, typically between \(54^\circ \text{C}\) and \(58^\circ \text{C}\).
Practical Applications in Heat Packs
The dramatic and rapid release of heat during crystallization is the physical principle utilized in reusable heat packs and hand warmers. These commercial products contain a sealed pouch filled with a supercooled, supersaturated solution of sodium acetate. The liquid within the pack is stable at room temperature but remains ready to crystallize upon activation.
Activation is achieved by flexing a small metal disc enclosed within the solution, which introduces the necessary mechanical shock or nucleation site. This action causes the entire volume of liquid to rapidly solidify in an exothermic chain reaction. The pack warms up almost instantly, providing a consistent source of heat for a period of time.
Once the heat pack has cooled and the sodium acetate is completely solid, it can be easily “recharged” for future use. The pack is simply placed in boiling water, which supplies the endothermic energy needed to melt the solid crystals and redissolve the salt. This process returns the solution to its supercooled liquid state, ready to be activated again once it cools down to room temperature.