When two aqueous solutions of ionic compounds are mixed together, the resulting reaction may produce a solid precipitate. This guide will show how to use the solubility rules for inorganic compounds to predict whether or not the product will remain in solution or form a precipitate.
Aqueous solutions of ionic compounds are comprised of the ions making up the compound dissociated in water. These solutions are represented in chemical equations in the form: AB(aq) where A is the cation and B is the anion.
When two aqueous solutions are mixed, the ions interact to form products.
AB(aq) + CD(aq) → products
This reaction is generally a double replacement reaction in the form:
AB(aq) + CD(aq) → AD + CB
The question remains, will AD or CB remain in solution or form a solid precipitate?
A precipitate will form if the resulting compound is insoluble in water. For example, a silver nitrate solution (AgNO3) is mixed with a solution of magnesium bromide (MgBr2). The balanced reaction would be:
2 AgNO3(aq) + MgBr2 → 2 AgBr(?) + Mg(NO3)2(?)
The state of the products needs to be determined. Are the products soluble in water?
According to the solubility rules, all silver salts are insoluble in water with the exception of silver nitrate, silver acetate and silver sulfate. Therefore, AgBr will precipitate out.
The other compound Mg(NO3)2 will remain in solution because all nitrates, (NO3)-, are soluble in water. The resulting balanced reaction would be:
2 AgNO3(aq) + MgBr2 → 2 AgBr(s) + Mg(NO3)2(aq)
Consider the reaction:
KCl(aq) + Pb(NO3)2(aq) → products
What would be the expected products and will a precipitate form?
The products should rearrange the ions to:
KCl(aq) + Pb(NO3)2(aq) → KNO3(?) + PbCl2(?)
After balancing the equation,
2 KCl(aq) + Pb(NO3)2(aq) → 2 KNO3(?) + PbCl2(?)
KNO3 will remain in solution since all nitrates are soluble in water. Chlorides are soluble in water with the exception of silver, lead and mercury. This means PbCl2 is insoluble and form a precipitate. The finished reaction is:
2 KCl(aq) + Pb(NO3)2(aq) → 2 KNO3(aq) + PbCl2(s)
The solubility rules are a useful guideline to predict whether a compound will dissolve or form a precipitate. There are many other factors that can affect solubility, but these rules are a good first step to determine the outcome of aqueous solution reactions.
Tips for Success Predicting a Precipitate
The key to predicting a precipitate is to learn the solubility rules. Pay particular attention to compounds listed as "slightly soluble" and remember that temperature affects solubility. For example, a solution of calcium chloride is typically considered soluble in water, yet if the water is cold enough, the salt doesn't readily dissolve. Transition metal compounds may form a precipitate under cold conditions, yet dissolve when it's warmer. Also, consider the presence of other ions in a solution. This can affect solubility in unexpected ways, sometimes causing a precipitate to form when you didn't expect it.
- Zumdahl, Steven S. (2005). Chemical Principles (5th ed.). New York: Houghton Mifflin. ISBN 0-618-37206-7.