In this tutorial, you will be introduced to the common ion effect. You will learn what it is, how it works and about its overall effect on solubility and precipitates in both chemical reactions and the real world.
Topics Covered in Other Articles
- Le Chatelier’s Principle
- Equilibrium Constant
- Lewis Acids & Bases
- Acid- Base Theories
- Buffer Solutions
- What is pH?
What is the Common Ion Effect?
First of all, what is a common ion? Let’s say there is a solution of zinc hydroxide, and zinc ions were added to the solution. Because it is found in both the solution and the salt that is added, the zinc ions would be the common ions in this example.
The common ion effect describes an ion’s effect on the solubility equilibrium of a substance. If a soluble compound consisting of a common ion is added, it can decrease the concentration of that ion within the solution; this can result in a change in the equilibrium point of the solution. This is seen when analyzing the solubility of weak electrolytes, such as salts.
This is a result of the concepts of Le Chatelier’s principle about ions’ ability to associate or dissociate. Read our tutorial on Le Chatelier’s principle to learn about it in more depth!
Effect on Solubility
Adding a Common Ion
If there is a solution containing a salt and a common ion (of the salt) is added, there will be precipitation of the salt. The concentration of the individual ions will decrease because they are being used to create the salt precipitate. Until the solubility equilibrium is reached, the concentration of the ions will continue to decrease.
Using the common ion definition example, this would be like adding zinc ions to the zinc hydroxide solution; the common ion would be zinc.
Adding Solution containing a Common Ion
When you add a solution containing a common ion, it has the same effect on solubility as just adding the common ions themselves. The incoming solution will dissociate, causing an increase in the salt precipitation, and a decrease in the ion concentration. The other ion in the incoming solution will usually not have a vast effect on the solubility equilibrium.
Using the common ion definition example, this would be like adding sodium hydroxide to the zinc hydroxide solution; the common ion would be hydroxide.
Common Ion Effect in Buffer Solutions
There is a change in pH when a conjugate ion is added to a buffer solution; the common ion effect is responsible for this change.
Let’s think about an acid buffer solution with a strong electrolyte (i.e. its conjugate base salt) added and dissolved; the electrolyte will dissociate, producing both common ions and ions that will partially ionize the acid. The generated common ions will counteract the ionization of the acid. Due to Le Chatelier’s Principle, the equilibrium will shift to favor the reactants (to the left). The pH increases because the ionization (and dissociation) of the acid is decreases.
Real-World Example of Common Ion Effect
When water is found underground, it is mixed with natural materials (i.e. rocks); therefore, it may contain limestone or chalk. Sodium carbonate, which is very soluble, can be added to decrease the hardness of water, or the levels of calcium and magnesium. This results in the creation of calcium carbonate, which is not very soluble.
More sodium carbonate is then added; in this case, carbonate ends up being the common ion. This addition results in the precipitation of calcium carbonate out of the water solution. Read about water treatment in depth here!
Example of the Common Ion Effect
Let’s take a solution of calcium sulfate as an example. In a saturated solution, calcium sulfate is in equilibrium with calcium ions and sulfate ions. Most of the calcium sulfate will remain as molecules, but a small percentage will dissociate into ions.
Let’s add some calcium chloride to this saturated solution. In this case, the concentration of the calcium ion in the solution would increase. The product of [Ca2+ ] times [SO42− ] would also increase, and would momentarily be greater than the Ksp. The equilibrium above would shift to the left because of the added calcium ion, according the Le Chatelier’s principle. More calcium sulfate would precipitate out of the solution until the ion product once again becomes equal to the Ksp. Notice that when the equilibrium is again reached, the concentrations of the calcium and sulfate ions would no longer be equal to each other – the calcium ion concentration would be much higher, but the product of the concentrations would still be equal to the Ksp.
This is a great example of the common ion effect. Adding calcium chloride to the saturated solution of calcium sulfate causes additional calcium sulfate to precipitate from the solution, lowering its effective solubility. Adding a compound or solution containing sulfate ions, such as sodium sulfate, would result in a similar effect.