Dissolving common table salt, or sodium chloride (\(\text{NaCl}\)), in water profoundly transforms the liquid’s fundamental characteristics. This interaction is not simple mixing, but a dynamic chemical process affecting water molecules at their most intimate level. The resulting solution, known as saline water or brine, possesses properties distinctly different from pure water, altering its thermal behavior, density, and electrical capacity. Understanding this change requires looking closely at how the polar nature of the water molecule engages with the ionic structure of salt.
The Molecular Interaction: Dissociation and Hydration
Water (\(\text{H}_2\text{O}\)) is a highly polar molecule, meaning it has a lopsided distribution of electrical charge. The oxygen atom holds a slight negative charge, while the two hydrogen atoms carry slight positive charges, creating a dipole moment. This inherent polarity makes water an effective solvent for substances held together by ionic bonds, such as sodium chloride.
When a salt crystal is introduced to water, the water molecules begin to interact with the crystal lattice. The positively charged hydrogen ends of the water molecules are attracted to the negatively charged chloride ions (\(\text{Cl}^-\)). Conversely, the negatively charged oxygen ends are drawn to the positively charged sodium ions (\(\text{Na}^+\)). This strong electrical attraction, called an ion-dipole interaction, overcomes the ionic bond holding the \(\text{Na}^+\) and \(\text{Cl}^-\) ions together in the solid crystal.
The salt crystal then undergoes dissociation, pulling the sodium and chloride ions completely apart and releasing them into the solution as free, charged particles. Once separated, each ion becomes surrounded by a shell of water molecules, known as a hydration shell. For example, the \(\text{Na}^+\) ion is encapsulated by water molecules oriented with their oxygen atoms facing inward.
The formation of these stable hydration shells isolates the ions, preventing them from recombining into solid salt. This process is thermodynamically favorable because the energy released by the strong ion-dipole bonds forming between the ions and water outweighs the energy required to break the ionic bonds in the salt crystal.
Impact on Thermal Properties
The presence of dissolved salt ions directly interferes with the natural phase transitions of water, changing both its boiling and freezing points. These changes are classified as colligative properties, meaning they depend primarily on the concentration of dissolved particles, not their chemical identity.
When water is cooled, its molecules slow down and arrange themselves into the highly ordered, hexagonal crystalline structure of ice. The dissolved \(\text{Na}^+\) and \(\text{Cl}^-\) ions disrupt this organizational process. Surrounded by hydration shells, the ions act as physical impediments that prevent water molecules from locking into the rigid ice lattice at the normal freezing temperature of \(0^{\circ}\text{C}\).
To overcome this interference and force the solution to freeze, the temperature must drop even lower, a phenomenon known as freezing point depression. For instance, a concentrated salt solution can depress the freezing point of water to approximately \(-21^{\circ}\text{C}\). This mechanism is why salt is applied to roads in winter; it lowers the melting point of the ice, allowing it to turn back into a liquid even when the air temperature is below the freezing point of pure water.
Conversely, the presence of ions makes it more difficult for water to transition into the gaseous, or vapor, phase, causing boiling point elevation. In pure water, boiling occurs when the water’s vapor pressure equals the surrounding atmospheric pressure. In salt water, the attractive ion-dipole forces within the hydration shells hold the water molecules more tightly, making it harder for them to escape the liquid surface as steam.
More energy, and therefore a higher temperature, is required to break these stronger attractions and reach the necessary vapor pressure for boiling. Adding approximately 58 grams of salt to one kilogram of water can raise the boiling point by about \(0.5^{\circ}\text{C}\). While this effect is measurable, the small amount of salt typically added to cooking water results in a negligible increase in temperature.
Changes to Physical Characteristics
Beyond altering temperature-related phase changes, dissolved salt ions modify the macroscopic physical characteristics of water, specifically its density and electrical conductivity. Adding salt significantly increases the overall mass of the solution without causing a proportional increase in volume, as the ions fit neatly into the spaces between water molecules.
This addition of mass per unit volume directly raises the density of the solution. Pure water has a density near \(1.0\) gram per milliliter, but a highly saline solution, such as that found in the Dead Sea, can have a density around \(1.24\) grams per milliliter. The increased density explains why objects are more buoyant in salt water, as the denser liquid exerts a greater upward force.
The most dramatic physical change is the solution’s ability to conduct electricity. Pure \(\text{H}_2\text{O}\) is a poor electrical conductor because it contains very few charged particles capable of carrying a current. However, when \(\text{NaCl}\) dissolves, it releases a high concentration of free-moving, charged ions (\(\text{Na}^+\) and \(\text{Cl}^-\)).
These mobile ions allow the solution to become an excellent electrolyte. When an electrical voltage is applied, positive ions are drawn toward the negative electrode, and negative ions move toward the positive electrode, effectively transferring electrical charge through the water. The electrical conductivity increases proportionally with the concentration of dissolved salt, reflecting the number of charge carriers available.