Solubility Equilibria and the Solubility Product

Theory — Solubility Equilibria and the Solubility Product

When an ionic solid is placed in water, it dissolves until the solution is saturated and a dynamic equilibrium is reached between the solid and its dissolved ions.

1. The solubility product

For a salt that dissolves as MaXb → a M+ + b X-, the equilibrium constant for dissolving is the solubility product:

2. Molar solubility

The molar solubility s is the moles of salt that dissolve per litre. It links directly to K_sp. For a 1:1 salt such as AgCl, Ksp = s2, so s = √Ksp. For a 1:2 salt such as CaF₂:

3. The common-ion effect

Adding an ion the salt already contains shifts the equilibrium back toward the solid, so the salt dissolves less. Silver chloride, for example, is far less soluble in a sodium chloride solution than in pure water. This is Le Chatelier’s principle applied to solubility.

4. Predicting precipitation with Q

To decide whether mixing two solutions makes a precipitate, compute the ion product Q with the same form as K_sp but using the actual mixed concentrations, and compare:

Apparatus

Solubility work uses tools to form, separate, and measure precipitates and ion concentrations. In the simulation these are modelled, but the readings match what each instrument would give.

Instructions

Work through both tabs. Calculate first by hand, then press the button to compare.

Part A — Molar solubility from K_sp

• Choose a salt and read its K_sp and dissolving stoichiometry.
• Compute the molar solubility s (use s = √K_sp for a 1:1 salt, or s = (K_sp/4)^1/3 for a 1:2 salt).
• Enter your value and press Check; it compares within 4 percent.

Part B — Precipitation and the common ion

• Set the mixed ion concentrations, or add a common ion.
• Predict whether a precipitate forms by comparing Q with K_sp, then press Check.

Simulation

Team Questions

Example Report

Practice Questions