CHEM 102 Santa Monica College Determination of a Solubility Product Constant Lab

Determination of a Solubility Product ConstantCHEMISTRY 102
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Objectives
In this experiment you will determine the solubility of calcium iodate (Ca(IO3)2).
By determining the solubility of this slightly soluble ionic compound, a Ksp can be calculated.
Explanation
The equilibrium between solid calcium iodate and its free ions in a saturated solution is given by:
Ca(IO3)2(s)
Ca2+(aq) + 2 IO3– (aq)
Ksp = [Ca2+] [IO3–]2
If either the concentration of free Ca2+ or IO3– ions is known in the saturated solution, the solubility product
constant (Ksp) can be calculated. The IO3– ion concentration can be determined through an oxidation-reduction titration
with a standardized solution of sodium thiosulfate in the presence of iodide, using starch as an indicator.
Iodide (I–) reacts with iodate (IO3–) in acid solution to give molecular I2 and water.
5 I– (aq) + IO3– (aq) + 6 H+ (aq) ⎯⎯→ 3 I2 (aq) + 3 H2O(l)
(1)
Molecular I2 reacts with thiosulfate ions (S2O32–) forming iodide ion & tetrathionate ion (S4O62–).
I2 (aq) + 2 S2O32– (aq) ⎯⎯→ 2 I– (aq) + S4O62– (aq)
(2)
Starch is used as an indicator because it reacts with molecular iodine reversibly forming a blue color. Thus, when all
of the molecular iodine has been consumed the blue color is no longer present signaling the end point of the titration.
When these two equations are combined the net reaction for the titration is obtained.
5 I– (aq) + IO3– (aq) + 6 H+ (aq) ⎯⎯→ 3 I2 (aq) + 3 H2O(l) (1)
3 I2 (aq) + 6 S2O32– (aq) ⎯⎯→ 6 I– (aq) + 3 S4O62– (aq)
(2)
–
2–
+
–
2–
IO3 (aq) + 6 S2O3 (aq) + 6 H (aq) ⎯⎯→ I (aq) + 3 S4O6 (aq) + 3 H2O(l) (net eqn.)
Procedure (wear safety goggles)
Preparing saturated solutions of calcium iodate
1. Prepare calcium iodate by adding 50.0 mL of 0.200 M KIO3 to 10.0 mL of 1.00 M Ca(NO3)2.
2. Stir the mixture with a stirring rod, a white precipitate of calcium iodate should form.
3. Allow the mixture to stand for ~10 minutes. (while you’re waiting, prepare gravity filtration setup)
4. Filter the precipitate using gravity filtration.
5. Wash the precipitate on the filter paper with three small portions (~ 5 mL each) of distilled water.
6. Locate two clean beakers with capacities of at least 100 mL and label them A and B.
7. Put ~1/3 of your dry precipitate in beaker A & ~ 1/3 of your dry precipitate in beaker B.
(Save ~1/3 of precipitate in case of a mishap.)
8. Use a graduated cylinder to add 40.0 mL of distilled water to beaker A and 80.0 mL of distilled water to beaker B.
9. Stir contents of each beaker thoroughly, using a different stirring rod for each beaker.
10. Stir contents of both beakers every 5 minutes for 30 minutes (leave stirring rods in beakers between stirrings).
11. Filter the mixture from beaker A using an absolutely dry funnel and dry filter paper.
Catch the filtrate in a clean dry beaker (label beaker: filtrate A).
12. Filter the mixture from beaker B using an absolutely dry funnel and dry filter paper.
Catch the filtrate in a clean dry beaker (label beaker: filtrate B).
Analyzing the saturated solution of calcium iodate Obtain a buret and a 10.00-mL volumetric pipette from stock room.
1a. Rinse the 10.00-mL volumetric pipette with distilled water and shake out as much water as possible from the pipette.
1b. Rinse the pipette with two small portions (~ 2 mL each) of filtrate A. (discard these portions)
1c. Pipet precisely 10.00 mL of filtrate A into a clean 250-mL Erlenmeyer flask (label flask: filtrate A-1).
2. Prepare the buret for titration (see Chem 101 Acid-base Titrations). Fill the buret with sodium thiosulfate.
3a. Add ~20 mL of distilled water to the Erlenmeyer flask (filtrate A-1).
3b. Dissolve ~1 g of solid KI in the Erlenmeyer flask (filtrate A-1).
3c. Add 10 drops of 6 M HCl to the Erlenmeyer flask (filtrate A-1).
3d. Add 10 drops of 1.0 % starch solution to the Erlenmeyer flask (filtrate A-1).
4a. Record the molarity & initial volume of the sodium thiosulfate solution in the buret in the Experimental Results table.
(Record the volumes of the buret to the nearest 0.02 mL.)
4b. Start titrating; the end point is reached when the solution becomes colorless.
4c. Record the final volume of the sodium thiosulfate solution in the buret at the end point.
5. Repeat steps 1c to 4 with a second sample of filtrate A.
6. Repeat steps 1a to 4 with the first sample of filtrate B.
7. Repeat steps 1c to 4 with a second sample of filtrate B.
A. For each Titration trial
a) Calculate the moles of sodium thiosulfate used to reach the endpoint.
b) Use the stoichiometry of the net equation (on page 1) to calculate the moles of iodate consumed.
c) Calculate the molarity of the iodate ion in the original filtrate.
d) Box or highlight the molarity of the iodate ion in the original filtrate.
Titration trial 1
Titration trial 2
Titration trial 3
Titration trial 4
Experimental Results Table (Complete this table after completing the calculations for each trail)
Titration trial
1
2
3
A-1
A-2
B-1
Filtrate
0.0243 M
0.0243 M
0.0243 M
[Na2S2O3] (in buret)
36.74
36.96
36.88
Vfinal buret (sodium thiosulfate)
0.12 mL
0.08 mL
0.04 mL
Vinitial buret (sodium thiosulfate)
VTotal (sodium thiosulfate)
a) moles (sodium thiosulfate)
b) moles (iodate)
c) [IO3–] (in original filtrate)
B. Use the average value of [IO3–] to calculate a Ksp for Ca(IO3)2.
[IO3–]Average(in original filtrate) = _____________
Ksp of Ca(IO3)2 = ______________
C. Should the IO3– concentrations from filtrates A and B be the same or different? Explain why?
4
B-2
0.0243 M
36.86
0.14 mL