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2.12: Solubility Equilibria: Overview

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JoVE Core
Analytical Chemistry

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Solubility Equilibria: Overview
 
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2.12: Solubility Equilibria: Overview

When a substance such as sodium chloride is added to water, it dissolves, forming an aqueous solution. The extent of dissolution is called solubility. The process of dissolution can exist in equilibrium, just like other chemical processes. Solubility equilibria are also called precipitation equilibria because the process of solubility can be reversible. The reverse of the solubility process is called precipitation.

Solubility is important in biological and environmental processes. A notable example of biological solubility is the effect of foods on tooth enamel, which consists of the mineral hydroxyapatite. Eating foods rich in sugars produces organic acids that dissolve hydroxyapatite, leading to tooth decay. Another example is calcium oxalate, a sparingly soluble salt that, if not flushed out by drinking sufficient water, will precipitate and form kidney stones over time.

Solubility equilibria follow Le Chatelier's principle, which states that if any force is applied to a reaction at equilibrium, the net reaction shifts towards whichever direction helps mitigate the stress from this force. Factors affecting the solubility equilibria of a sparingly soluble salt include temperature, solvent, common ion effect, pH of the solution, and the extent of complex ion formation. 

When a sparingly or moderately soluble salt is added to a solvent or solution, a dynamic equilibrium exists between the dissociated ions and the solid compound in the saturated solution. The equilibrium constant of a sparingly soluble salt is the solubility product constant or solubility product, Ksp, which is independent of the concentration of the solid form of salt because the concentration of the solid in a saturated solution is constant.

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Keywords: Solubility Equilibria Precipitation Equilibria Solubility Dissolution Aqueous Solution Sparingly Soluble Salt Solubility Product Constant Ksp Le Chatelier's Principle Temperature Solvent Common Ion Effect PH Complex Ion Formation

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