Chapter 16 Worksheet 3 (ws16.3)
Quantitative Analysis of Solubility Equilibria
Solubility of Ionic Compunds
In chapter 4, some solubility guidelines for ionic compounds were given in Table 4.2. A similar table from another textbook is given below. These compounds are classified as either soluble or insoluble. In reality many ionic compounds classified as insoluble are slightly soluble. We will now develop the tools to handle solubility quantitatively using equilibrium constants for dissolution reactions. Such an equilibrium constant is called Ksp where sp stands for solubility product. Some Ksp values are listed in table 16.2 in your textbook and reproduced at the end of this worksheet.
We are interested in calculating the solubility of slightly soluble salts such as silver carbonate.
Ksp is called the “solubility product” but it is NOT the solubility of the compound. It is the equilibrium constant for the dissolving (dissolution) reaction. For example:
Ag2CO3 (s) = 2Ag+(aq) + CO32−(aq) Ksp = [Ag+]2[ CO32−] = 8.1 x 10−12
When the reaction is at equilibrium, the solution is saturated.
The solubility of an ionic compound is its concentration in a saturated solution (its equilibrium concentration). This is not quite right since after an ionic compound dissolves, 100% of it dissociates. In the reaction above, the solubility of silver carbonate is actually the equilibrium concentration of the carbonate ion or ½ of the silver ion concentration.
In problems 1 and 2, you are given the solubility and asked to calculate Ksp. In problem 3 you are asked to do the opposite (given Ksp calculate solubility).
1. a. Write the balanced equation for calcium hydroxide dissolving in H2O.
b. The solubility of calcium hydroxide is 0.012 M. What are the concentrations of calcium ion and hydroxide ion in a saturated solution?
[Ca2+]eq =
[OH-]eq =
c. Write the equilibrium constant expression and calculate Ksp for this reaction.
2. a. Write the balanced equation for calcium phosphate dissolving in H2O.
b. The solubility of calcium phosphate is 2.21 x 10-4 g/L. What are the molar concentrations of calcium ion and phosphate ion in a saturated solution? (molecular wt. of calcium phosphate = 310.18 g/mole)
[Ca2+]eq =
[PO43-]eq =
c. Write the equilibrium constant expression and calculate Ksp for this reaction.
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3. Calculate the molar solubilities of the following:
a. CuCO3 (Ksp = 2.5 x 10-10)
dissolution reaction:
I
C
E
Solubility = ______M
b. Ag2SO4 (Ksp = 1.7 x 10-5)
dissolution reaction:
I
C
E
Solubility = ______M
4. In a saturated solution of magnesium phosphate, [Mg2+] = 2.4 x 10-6 M and [PO43-] = 1.6 x 10-6 M. What is the solubility of magnesium phosphate?
Will a precipitate form?
5. In chapter 4, you studied precipitation reactions in which two solutions of soluble salts were mixed to produce at least one insoluble salt.
a. Use table 5.1 to predict whether or not a precipitate will form when a solution of lead (II) nitrate is mixed with a solution of sodium chloride.
b. Write the molecular, total ionic, and net ionic equations for the reaction.
Molecular:
Total ionic:
Net ionic:
c. In this type of problem one of the reactants is often a sodium salt and the other reactant is often a nitrate. Why? (Consult table 5.1).
6. In the previous problem, it was assumed that a precipitate would form because PbCl2 is classified as “insoluble”. Although the solubility of PbCl2 is very small, it is not zero. In general, a reaction between two soluble salts produces a precipitate only when the amount of product formed exceeds the solubility of the compound. To predict whether or not a precipitate will form, simply compare Q to Ksp:
If Q < Ksp, no precipitate will form
If Q > Ksp, a precipitate will form
If Q = Ksp, the solution is saturated (no precipitate will form)
a. Look up Ksp for PbCl2 in table 16.2 and calculate its solubility.
b. If 5.00 mL of 0.300 M lead (II) nitrate is mixed with 10.0 mL of 0.150 M sodium chloride, a precipitate of PbCl2 forms as predicted above. However, if 5.00 mL of 0.0300 M lead (II) nitrate is mixed with 10.0 mL of 0.0150 M sodium chloride, no precipitate forms. Explain these observations by calculating Q for each reaction and compare Q to Ksp PbCl2. The calculations for the more concentrated solutions are done for you. Do the same calculations for the less concentrated solutions.
Calculations for the more concentrated solutions
Since two solutions are mixed, you must calculate the initial concentration of each salt using the dilution equation: M1V1 = M2V2 (M1 and V1 are the molarity and volume before dilution. M2 and V2 are the molarity and volume after dilution.)
[Pb2+] =
[Cl-] =
PbCl2 (s) ⇌Pb2+ (aq) + 2Cl- (aq)
Ksp = [Pb2+][Cl-]2 = 1.7 x 10-5
(Remember: Ksp is the equilibrium constant for the dissolving reaction, not the precipitation reaction!)
Q = (0.100)(0.100)2 = 1.00 x 10-3
Q > Ksp so the dissolving reaction will proceed in the reverse direction and a precipitate will form.
Calcuations for the less concentrated solutions
7. Will a precipitate form if 90.0 mL of 0.200 M tin (II) nitrate is added to 10.0 mL of 0.200 M sodium iodide? (Ksp for tin iodide is 1.0 x 10-4)
Common-Ion Effect: Solubility of ionic compounds
Note: The “common-ion effect” is nothing new! It is simply a name for what happens when you add a product to a reaction at equilibrium. (It shifts to the left.) In this case the added product happens to be an ion!
8. a. Calculate the molar solubility of PbI2 (Ksp = 5.0 x 10-7) in water.
dissolution reaction:
I
C
E
Solubility = ______M
b. Calculate the molar solubility of PbI2 (Ksp = 5.0 x 10-7) in 0.20 M KI.
dissolution reaction:
I
C
E
Solubility = ______
The effect of pH on the solubility of ionic compounds
9. For many compounds, such as Ca(OH)2 or CaCO3, solubility will vary with pH of the solution. This is because the anions, OH- and CO32-, are bases. Using LeChatelier’s principle, predict how the solubility of Ca(OH)2 will vary with each of the following changes.
Ca(OH)2 (s) ⇄ Ca2+ (aq) + 2 OH- (aq)
a. Lowering the pH of the solution by adding a strong acid, HCl.
b. Raising the pH of the solution by adding a strong base, NaOH.
7. Circle the compounds that are more soluble in acidic solution than in pure water.
Al(OH)3, NaCl, BaF2, CuBr, NaClO3, NaClO2, NaNO3, NaNO2
8. Calculate the solubility of barium hydroxide (Ksp = 5.0 x 10-3) in pure water and in a buffer whose pH = 13 (For pH 13, it is a common ion problem. The common ion is hydroxide ion so you need to calculate the initial hydroxide ion concentration and set up an ICE table.)
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