# Insoluble Salt Precipitation

Learning Goals

You will learn how to use the reaction quotient to predict the state of saturation of an insoluble salt water solution, and the concentration of ions needed to initiate the precipitation of the salt in an insoluble salt solution.

Synopsis

In Section 16.3 you learned about the reaction quotient and applied it to equilibrium reactions. Here we will use the reaction quotient to determine if a solution of a set of ions will produce a precipitate. Again we will use our hypothetical salt equilibrium AaBb aA+ + bB-. The reaction quotient is written in exactly the same way as in Chapter 16: Q = [A+]a[B-]b. Remember that the difference between Q and Ksp is that for Ksp we use the equilibrium concentration of products. The same equilibrium condition used in Chapter 16 apply here [another example of the building-block nature of chemistry], but the interpretations are for a salt in solution. The conditions are:

• If Q = Ksp the reaction is at equilibrium. In this application this means that the dissociation of the salt is at equilibrium with its ions in solution so the solution is saturated. No more salt will dissolve.
• If Q > Ksp the reaction is proceeding toward the reactant side. In this application this means that the solution is supersaturated and no more salt will dissolve in solution.
• If Q less than Ksp the reaction is proceeding toward the product side. In this application this means that the solution is not saturated and still more of the salt will dissolve in the solution. However, there is another way to look at this situation. If we actually started with ions in solution instead of the solid salt, Q less than Ksp means that we can still add more ions in solution before precipitation of the salt occurs (when Q will become equal to Ksp).
So we can now work two very useful types of problems:

(1) given the ion concentrations in solution, is the dissolution process at equilibrium, will the salt continue to dissolve, or will more precipitate form, and

(2) what is the ion concentration(s) that will cause a precipitate to form. Remember in the second application that we still write the reaction as AaBb aA+ + bB- even though the temptation is to reverse the reaction.

Review Questions

1. AuCl is placed in a beaker with water and the concentrations of Au+ and Cl- ions are each determined to be 2.84 x 10-6 M. What is the equilibrium status of the solution?
2. If the above concentration of ions had been introduced into the water other than by the addition of AuCl, what concentration of ions could be added so that precipitation of AuCl would first begin to occur?
3. If NaCl where placed in the above solution so that the Cl- concentration were 0.0020 M, would any Au+ ions remain in solution?

Web Author: Dr. Leon L. Combs