Sodium chloride (\(\text{NaCl}\)), commonly known as table salt, is extremely soluble in water. This high solubility is a direct consequence of the intrinsic electrical properties of both the salt and the water molecules. Water is often called the “universal solvent” because its unique molecular structure allows it to dissolve a wide range of compounds. The process of dissolution is a physical event driven by an energy exchange that separates the salt’s crystal structure into individual components.
The Polarity of Water and Ionic Bonds in Salt
The ability of water to dissolve salt begins with the fundamental structure of both compounds. Sodium chloride is an ionic compound held together by strong electrostatic attraction, known as an ionic bond, between two oppositely charged ions. Within its solid crystal lattice, positively charged sodium ions (\(\text{Na}^+\)) and negatively charged chloride ions (\(\text{Cl}^-\)) are arranged in a repeating structure.
These ionic bonds are stable and require significant energy to break apart. Water (\(\text{H}_2\text{O}\)) possesses polarity, which is the key to overcoming this stability. A water molecule is bent in shape, giving it a partial negative charge near the oxygen atom and partial positive charges near the two hydrogen atoms. This uneven distribution of electrical charge creates a strong dipole moment, making one end of the molecule slightly negative and the other end slightly positive.
How Sodium Chloride Dissolves in Water
When salt comes into contact with water, the polar water molecules use their electrical charges to interact with the crystal lattice. The partially negative oxygen end of the water molecule is strongly attracted to the positive sodium ions in the salt crystal. Simultaneously, the partially positive hydrogen ends of the water molecule are drawn toward the negative chloride ions. This attraction between the water’s dipole and the salt’s ions is called an ion-dipole interaction.
As water molecules cluster around the salt crystal’s surface, their collective pulling force overcomes the strong ionic bonds holding the ions together. This process requires an input of energy, known as the lattice energy, to separate the ions from the crystal structure. Once freed, the individual ions are surrounded by a shell of water molecules, known as a hydration shell or solvation shell.
The formation of stable ion-dipole attractions releases energy called hydration energy. For sodium chloride, the energy released through hydration is nearly equal to the energy required to break the crystal lattice apart. This energy balance explains the ease of dissolution, as the solution process is slightly endothermic, meaning it absorbs a small amount of heat from the surrounding water. Once hydrated, the water molecules shield the \(\text{Na}^+\) and \(\text{Cl}^-\) ions, preventing them from recombining into solid salt.
Temperature’s Effect on Sodium Chloride Solubility
For most solid substances, increasing the water temperature significantly increases the amount that can dissolve. However, sodium chloride exhibits unusual behavior: its solubility changes very little across a wide temperature range. At \(20^\circ\text{C}\), approximately 36.0 grams of \(\text{NaCl}\) can dissolve in 100 grams of water.
Even when the water is heated to \(100^\circ\text{C}\), the solubility only increases slightly to about 39.2 grams per 100 grams of water. This minimal change is due to the near-balance between the energy needed to break the lattice and the energy released by hydration. Because the dissolution of \(\text{NaCl}\) is nearly thermoneutral, adding heat does not dramatically shift the equilibrium toward more dissolving. This means that \(\text{NaCl}\) is prone to quickly forming a saturated solution, which is the state where the maximum possible amount of solute has been dissolved at a given temperature.