Experiment 22: Molar Solubility (Part A) and Common Ion Effect (Part B)
Background Information
A saturated solution of a salt with limited solubility has a reversible reaction and an
equilibrium expression. The equilibrium constant (Ksp) is called the solubility product.
Ag2SO4 ⇌ 2Ag+ + SO42- leads to Ksp = [Ag+]2[SO42-]
and Ca(OH)2 ⇌ Ca2+ + 2OH- leads to Ksp = [Ca2+] [OH-]2
The aqueous concentration of an aqueous compound (not its individual ions) is called the
molar solubility. If there is only one of an ion in the compound’s formula unit, the ion’s
concentration is the same as the molar solubility. Otherwise, the ion’s concentration is a
multiple of the molar solubility. That is, we can make equations for molar solubilities.
[Ag2SO4(aq)] = [SO42-] = (1/2) [ Ag+] and [Ca(OH)2(aq)] = [Ca2+] = (1/2)[OH-]
[OH-] is easy to determine with a titration against [H3O+], and is twice the [Ca(OH)2(aq)]
molar solubility, as seen in the equation above. Substituting the same molar solubility
value for [Ca2+] can provide the solubility product.
Ksp = [Ca2+] [OH-]2 = (1/2) [OH-][OH-]2 = (1/2)[OH-]3
If [Ca2+] is increased by adding CaCl2, then LeChatelier’s principle tells us that the
equilibrium will shift to the left. This equilibrium shift is called the common ion effect.
Now [Ca2+] is no longer the same as the molsol, [Ca(OH)2(aq)] or (1/2)[OH-]. Also, [Ca2+]
can be large enough that the shift is negligible in comparison to the amount from the
added CaCl2. Therefore, we would not calculate Ksp without knowing [Ca2+], the
concentration of the added common ion. However, molar solubility can still be
determined with a titration because it can still be calculated as [Ca(OH)2(aq)] = (1/2)[OH-].
As a result, we can experimentally determine the effect of the common ion on the molar
solubility, as predicted by LeChatelier’s principle.
(Also, see the Common Ion Effect section in the Ch 16 notes.)
Procedure Notes
1. Skip steps 1 and 2 of part A, and step 1 of part B. Using beakers, obtain 100 ml of
each solution from the supply bench. This quantity will be used to pipet 25-ml
samples for each set of three trials.
2. Review technique 16 (pages 29-33) for instruction on pipetting and titrating.
3. Skip the standard deviations on your report sheet.
Pre-Lab Hints:
1. a) Molar solubility is the saturated concentration of PbI2.
Use stoichiometry to the ratio of PbI2 to I-.
b) Put concentrations of Pb+2 and I- into the Ksp expression.
c) Refer to definition of common ion effect (pp 257-8).
Also, you may note that the calculated Ksp for this problem does not match
the value of Ksp in your textbook (Table 17.1, pg 703).
2. Consider how remaining DI water will affect the solution.
3. See footnote on page 259.
4. See step 2 of procedure 16A at the top of page 30.
5. Refer to discussion of molar solubility and common ion effect on pages 257-8.
6. (moles H3O+ added) = (moles OH- in solution)
[OH-] = (moles OH-) / (sample volume)
Review molar solubility and Ksp equations for [Ca(OH)2(aq)].
7. Compare endpoint pH (which will be <7) with indicator color change pH’s.
You may also want to look in Ebbing and Gammon at Figures 15.10 and 15.11
(pg 642), as well as Fig 16.14 (pg 686).
Lab Questions
1. Generally, some solid Ca(OH)2 is present in the sample. Given that acid removes
OH-1 from the right side of the dissociation reaction, what will happen to the solid and
its equilibrium when the acid is added? Explain how this will affect the moles of acid
needed to reach the stoichiometric point.
2. During the titration, you added (neutral) DI water from the squirt bottle to wash the
acid from the walls of the beaker. Explain whether or not this added water will affect
the moles of acid needed to reach the stoichiometric point.
How is the calculated Ksp affected?
3. It is not uncommon to add excess acid near the stoichiometric point, so that a larger
stoichiometric volume of acid solution is observed in the experiment. Refer to the
color pictures of pH indicators at the end of chapter 16 in your textbook. Explain
what happens to the color of the solution as the stoichiometric point is passed. Also,
explain what will happen to the experimental molar solubility and Ksp if the larger
volume is recorded.
4. Suppose the solution in part A was made with tap water instead of DI water.
Tap water typically contains what common ion that could affect your results?
Explain how the tap water could shift the solubility equilibrium
and change your experimental molar solubility.
5. Explain how you could use the experimental [OH-] in part A to find [Ca2+].
Also, explain whether [Ca2+] in part B would be larger or smaller than in part A.