By Natalie Andrews – Updated March 24, 2022
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At room temperature (about 25 °C), a standard 100 g (≈100 mL) of water can dissolve roughly 35 g of sodium chloride (table salt). This value represents the solubility limit; any additional salt will remain on the bottom of the vessel, indicating a saturated solution.
Solubility of most solids increases with temperature, though the magnitude varies by compound. For sodium chloride, the rise is modest: at 100 °C, the solubility climbs to about 39 g per 100 g water, allowing a few grams more salt to dissolve before saturation is reached.
When water is near its boiling point, it can accommodate up to ~40 g of salt before saturation, which explains why chefs often add a pinch of salt to boiling water to accelerate cooking.
In colder conditions, salt dissolves more slowly, and the presence of salt lowers the freezing point of water—a phenomenon known as freezing point depression. Adding salt to ice effectively prevents the ice from melting, a principle exploited in de‑icing roads.
An unsaturated solution contains more available solvent than solute; salt crystals are fully hydrated and dissolve completely. A saturated solution has reached equilibrium: the rate of dissolution equals the rate of crystallization, and any extra salt will simply settle out.
Below the normal freezing point—around –6 °C (–5.98 °F)—water can no longer accept additional salt molecules. The mixture then consists of solid ice crystals interspersed with salt grains.
Not all salts behave identically. Rock salt, which contains impurities, dissolves more slowly than pure table or canner’s salt. The presence of minerals or additives can interfere with the hydration process, extending the time required for full dissolution.
