1. Ionic Equilibria III: The Solubility Product Principle17
Ionic Equilibria III:
The Solubility Product Principle

2. Determination of Solubility Product Constants
Example 17-7: One liter of saturated silver chloride solution contains g of dissolved AgCl at 25oC. Calculate the molar solubility of, and Ksp for, AgCl.
The molar solubility can be easily calculated from the data:

3. Determination of Solubility Product Constants
Example 17-8: One liter of saturated calcium fluoride solution contains gram of CaF2 at 25oC. Calculate the molar solubility of, and Ksp for, CaF2.
Calculate the molar solubility of CaF2.

4. Determination of Solubility Product Constants
From the molar solubility, we can find the ion concentrations in saturated CaF2. Then use those values to calculate the Ksp.
Note: You are most likely to leave out the factor of 2 for the concentration of the fluoride ion!

5. Uses of Solubility Product Constants
Example 17-9: Calculate the molar solubility of barium sulfate, BaSO4, in pure water and the concentration of barium and sulfate ions in saturated barium sulfate at 25oC. For barium sulfate, Ksp= 1.1 x

6. Uses of Solubility Product Constants

7. Uses of Solubility Product Constants
Make the algebraic substitution of x’s into solubility product expression and solve for x, giving the ion concentrations.

8. Uses of Solubility Product Constants
Finally, to calculate the mass of BaSO4 in 1.00 L of saturated solution, use the definition of molarity.

9. Uses of Solubility Product Constants
Example 17-10: The solubility product constant for magnesium hydroxide, Mg(OH)2, is 1.5 x Calculate the molar solubility of magnesium hydroxide and the pH of a saturated magnesium hydroxide solution at 25oC.

10. Uses of Solubility Product Constants
Be careful, do not forget the stoichiometric coefficient of 2!

11. Uses of Solubility Product Constants
Substitute the algebraic expressions into the solubility product expression.

12. Uses of Solubility Product Constants
Solve for the pOH and pH.

13. The Common Ion Effect in Solubility Calculations
Example 17-11: Calculate the molar solubility of barium sulfate, BaSO4, in M sodium sulfate, Na2SO4, solution at 25oC. Compare this to the solubility of BaSO4 in pure water. (Example 20-3). (What is the common ion? How was a common ion problem solved in Chapter 19?)

14. The Common Ion Effect in Solubility Calculations
Write equations to represent the equilibria.

15. The Common Ion Effect in Solubility Calculations
Substitute the algebraic representations of the concentrations into the Ksp expression and solve for x.

16. The Reaction Quotient in Precipitation Reactions
Example 17-12: We mix 100 mL of M potassium sulfate, K2SO4, and 100 mL of 0.10 M lead (II) nitrate, Pb(NO3)2 solutions. Will a precipitate form?

17. The Reaction Quotient in Precipitation Reactions
Write out the solubility expressions.

18. The Reaction Quotient in Precipitation Reactions
Calculate the Qsp for PbSO4.
Assume that the solution volumes are additive.
Concentrations of the important ions are:

19. The Reaction Quotient in Precipitation Reactions
Finally, calculate Qsp for PbSO4 and compare it to the Ksp.

20. The Reaction Quotient in Precipitation Reactions
Example 17-13: What concentration of sulfide ions, from a soluble compound such as Na2S, is required to reduce the Hg2+ concentration to 1.0 x 10-8 M? For HgS, Ksp=3.0 x

21. The Reaction Quotient in Precipitation Reactions
Example 17-14: Refer to example What volume of the solution (1.0 x 10-8 M Hg2+ ) contains 1.0 g of mercury?

22. Fractional Precipitation
Example 17-15: If solid sodium chloride is slowly added to a solution that is M each in Cu+, Ag+, and Au+ ions, which compound precipitates first? Calculate the concentration of Cl- required to initiate precipitation of each of these metal chlorides.

23. Fractional Precipitation

24. Fractional Precipitation
Repeat the calculation for silver chloride.

25. Fractional Precipitation
Finally, for copper (I) chloride to precipitate.

26. Fractional Precipitation
Example 17-16: Calculate the percentage of Au+ ions that precipitate before AgCl begins to precipitate.
Use the [Cl-] from Example to determine the [Au+] remaining in solution just before AgCl begins to precipitate.

27. Fractional Precipitation
The percent of Au+ ions unprecipitated just before AgCl precipitates is:
Therefore, 99.9% of the Au+ ions precipitates before AgCl begins to precipitate.

28. Fractional Precipitation
A similar calculation for the concentration of Ag+ ions unprecipitated before CuCl begins to precipitate is:

29. Fractional Precipitation
The percent of Ag+ ions unprecipitated just before AgCl precipitates is:
Thus, % of the Ag+ ions precipitates before CuCl begins to precipitate.

30. Simultaneous Equilibria Involving Slightly Soluble Compounds
Example 17-17: If 0.10 mole of ammonia and mole of magnesium nitrate, Mg(NO3)2, are added to enough water to make one liter of solution, will magnesium hydroxide precipitate from the solution?
For Mg(OH)2, Ksp = 1.5 x Kb for NH3 = 1.8 x 10-5.
Calculate Qsp for Mg(OH)2 and compare it to Ksp.
Mg(NO3)2 is a soluble ionic compound so [Mg2+] = M.
Aqueous ammonia is a weak base that we can calculate [OH-].

31. Simultaneous Equilibria Involving Slightly Soluble Compounds

32. Simultaneous Equilibria Involving Slightly Soluble Compounds
Once the concentrations of both the magnesium and hydroxide ions are determined, the Qsp can be calculated and compared to the Ksp.

33. Simultaneous Equilibria Involving Slightly Soluble Compounds
Example 17-18: How many moles of solid ammonium chloride, NH4Cl, must be used to prevent precipitation of Mg(OH)2 in one liter of solution that is 0.10 M in aqueous ammonia and M in magnesium nitrate, Mg(NO3)2 ? (Note the similarity between this problem and Example )
Calculate the maximum [OH-] that can exist in a solution that is M in Mg2+.

34. Simultaneous Equilibria Involving Slightly Soluble Compounds

35. Simultaneous Equilibria Involving Slightly Soluble Compounds
Using the maximum [OH-] that can exist in solution, determine the number of moles of NH4Cl required to buffer 0.10 M aqueous ammonia so that the [OH-] does not exceed 3.9 x 10-5 M.

36. Simultaneous Equilibria Involving Slightly Soluble Compounds

37. Simultaneous Equilibria Involving Slightly Soluble Compounds

38. Simultaneous Equilibria Involving Slightly Soluble Compounds
Check these values by calculating Qsp for Mg(OH)2.

39. Simultaneous Equilibria Involving Slightly Soluble Compounds
Use the ion product for water to calculate the [H+] and the pH of the solution.

40. Complex Ion Equilibria
Example 17-19: How many moles of ammonia must be added to 2.00 L of water so that it will just dissolve mole of silver chloride, AgCl?
The reaction of interest is:

41. Complex Ion Equilibria
Two equilibria are involved when silver chloride dissolves in aqueous ammonia.

42. Complex Ion Equilibria
The [Ag+] in the solution must satisfy both equilibrium constant expressions. Because the [Cl-] is known, the equilibrium concentration of Ag+ can be calculated from Ksp for AgCl.

43. Complex Ion Equilibria

44. Complex Ion Equilibria
Substitute the maximum [Ag+] into the dissociation constant expression for [Ag(NH3)2]+ and solve for the equilibrium concentration of NH3.

45. Complex Ion Equilibria
The amount just calculated is the equilibrium concentration of NH3 in the solution. But the total concentration of NH3 is the equilibrium amount plus the amount used in the complex formation.

46. Complex Ion Equilibria
Finally, calculate the total number of moles of ammonia necessary.

47. Synthesis Question
Most kidney stones are made of calcium oxalate, Ca(O2CCO2). Patients who have their first kidney stones are given an extremely simple solution to stop further stone formation. They are told to drink six to eight glasses of water a day. How does this stop kidney stone formation?

48. Synthesis Question