Water, a fundamental substance, is commonly known to freeze at a specific temperature, transitioning from its liquid state to solid ice. Under standard atmospheric pressure, pure water consistently undergoes this phase change at 0°C (32°F). However, the presence of dissolved substances in water can significantly alter its physical properties, including its freezing point.
The Lowering of Freezing Point
Saltwater freezes at a temperature below the 0°C mark characteristic of pure water. The precise freezing point is directly influenced by the concentration of salt, referred to as salinity. As the amount of salt increases, the freezing point of the solution decreases further. For instance, average ocean water, which typically has a salinity of about 3.5%, begins to freeze at approximately -1.9°C to -2.0°C (28.4°F). This means that oceans and seas can remain in a liquid state even when temperatures drop below the freezing point of freshwater.
The Science Behind Freezing Point Depression
This phenomenon, where a dissolved substance lowers a liquid’s freezing point, is known as freezing point depression. It is a colligative property, meaning it depends on the number of solute particles, not their chemical identity.
When water freezes, its molecules typically arrange into a highly organized, repeating crystal structure, forming solid ice. Dissolved salt particles, such as sodium and chloride ions from table salt, interfere with this arrangement. These ions disrupt the hydrogen bonds that normally form between water molecules, which are necessary for water molecules to align into the rigid ice lattice. Consequently, a lower temperature is required for water molecules to overcome this disruption and solidify.
Real-World Examples of Saltwater Freezing
Freezing point depression is observed in natural and human-applied contexts. A notable natural example is sea ice formation in polar regions. While ocean water freezes below 0°C, the ice formed is considerably less salty than the surrounding seawater. This occurs through brine rejection, where salt ions are largely excluded from the ice crystal structure as water freezes, increasing the salinity of the remaining liquid water.
Another common application is using salt to de-ice roads and sidewalks during winter months. When salt, such as sodium chloride or calcium chloride, is applied to icy surfaces, it dissolves in any available liquid water, forming a saltwater solution. This solution has a lower freezing point than pure water, helping to melt existing ice or prevent new ice from forming. The effectiveness of de-icing salts can vary with temperature and salt type.