Solubility Product, Common Ion Effect and Selective Precipitation

The constant and its relation to solubility

For a sparingly soluble salt dissolving to give ions, the solubility product is the product of ion concentrations at saturation, each raised to its stoichiometric power.

Predicting precipitation

Compute the ionic product from the actual concentrations present and compare it with the solubility product:

Remember to account for dilution when two solutions are mixed. Using the original concentrations rather than those after mixing is a very common arithmetic error.

The common ion effect

Adding an ion already present in the equilibrium shifts it toward the undissolved solid, reducing solubility. This is simply Le Chatelier applied to a dissolution equilibrium.

It is used deliberately in gravimetric analysis: washing a precipitate with a solution containing one of its own ions, rather than with pure water, minimises the loss from dissolution during washing.

Selective precipitation

Where two ions both precipitate with the same reagent but have different solubility products, controlled addition of that reagent precipitates the less soluble one first.

1. Compute the reagent concentration needed to begin precipitating each ion.
2. The one requiring the lower concentration precipitates first.
3. Adding reagent up to just below the second threshold precipitates the first ion largely and leaves the second in solution.

The separation is good when the two solubility products differ by several orders of magnitude, and poor when they are close. This is exactly the principle used in the sulphide group separations of qualitative analysis, where pH controls the sulphide concentration.

The salt effect

Adding an inert electrolyte — one with no ion in common — increases solubility slightly, which looks like the opposite of the common ion effect.

The explanation is activity. Raising the ionic strength lowers the activity coefficients of the dissolving ions, so a higher concentration is required to reach the same activity product. Being able to explain this apparent contradiction distinguishes a thorough answer, and it connects the topic to Debye–Hückel theory.

Complications worth acknowledging

• Hydrolysis of an anion from a weak acid consumes it, so solubility depends on pH. Salts of weak acids dissolve more readily in acid.
• Complex formation can dissolve a precipitate entirely, which is why some precipitates redissolve in excess reagent.
• Very insoluble salts may not reach equilibrium quickly, so measured solubility can be kinetically limited.

The redissolution point is practically important: adding excess of a precipitating reagent does not always give more precipitate, and may give less.

Frequently asked questions

Why can solubility products not be compared directly?

Because the relationship between the constant and the solubility depends on stoichiometry. Only actual solubilities are comparable across different formula types.

Why does an inert salt increase solubility?

Because it raises the ionic strength, lowering activity coefficients, so higher concentrations are needed to reach the same activity product.

Why does a precipitate sometimes dissolve in excess reagent?

Because a soluble complex forms with the excess reagent, removing the ion from the dissolution equilibrium and driving the solid into solution.

How is pH used to separate sulphides?

By controlling the sulphide ion concentration. In acid it is very low, so only the least soluble sulphides precipitate; in base it is much higher, so the rest follow.