Glossary
Ion product (Qsp)
A value calculated using the current concentrations of ions in a solution, which can be compared to Ksp to predict whether a precipitate will form or if the solution is saturated, unsaturated, or supersaturated.
Example:
If you mix two solutions and calculate a Qsp that is greater than the Ksp, you know a precipitate will form.
Molar solubility
The concentration of the dissolved ionic compound in a saturated solution, expressed in moles per liter (mol/L or M). It represents the moles of solute that dissolve per liter of solution.
Example:
If the molar solubility of CaF2 is M, it means moles of CaF2 dissolve in every liter of water.
Precipitation
The process by which a solid forms from a solution, typically when the concentration of dissolved ions exceeds the solubility limit (Qsp > Ksp).
Example:
Mixing solutions of silver nitrate and sodium chloride causes precipitation of solid silver chloride.
RICE table
An acronym (Reaction, Initial, Change, Equilibrium) used to organize and solve equilibrium problems, particularly useful for calculating molar solubility or equilibrium concentrations.
Example:
When calculating the molar solubility of a sparingly soluble salt, setting up a RICE table helps track the changes in ion concentrations.
Saturated
A solution that contains the maximum amount of dissolved solute at a given temperature, where the dissolution and precipitation rates are equal, establishing equilibrium.
Example:
When sugar no longer dissolves in your tea and settles at the bottom, the solution is saturated.
Solubility
The maximum amount of a substance that can dissolve in a given amount of solvent at a specific temperature to form a saturated solution. It is an equilibrium process where a solid dissolves into its ions.
Example:
Sugar has high solubility in water, allowing you to dissolve many spoonfuls before it stops dissolving.
Solubility product constant (Ksp)
An equilibrium constant that describes the extent to which a sparingly soluble ionic compound dissolves in water. It is the product of the ion concentrations raised to their stoichiometric coefficients in a saturated solution.
Example:
A very small Ksp for AgCl () means very little silver chloride will dissolve in water.
Supersaturated
A solution that contains more dissolved solute than a saturated solution at the same temperature, making it unstable and prone to precipitation.
Example:
A solution of sodium acetate that has been carefully cooled can be supersaturated, and adding a tiny crystal will cause rapid crystallization.
Unsaturated
A solution that contains less dissolved solute than the maximum amount that can be dissolved at a given temperature, meaning more solute can still dissolve.
Example:
Adding a small amount of salt to water creates an unsaturated solution, as more salt can still dissolve.