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Solubility Product Constant of Calcium Iodate
Purpose: To determine titrimetrically the solubility product constant Ksp for the relatively insoluble salt Ca(IO3)2
Introduction
Whenever a sparingly soluble salt is added to water, an equilibrium is established between the undissolved salt and a saturated solution of the salt. The saturated solution contains the ions of the salt. For example, when lead (II) iodide is added to water, the equilibrium
+2 - PbI2(s) Pb (aq) + 2I (aq) (eq. 1) is established. Since this is an equilibrium process, the rate of dissolution of the solid to produce ions is equal to the rate of recombination, or precipitation, of the ions to form more solid. We can thus write an expression for the equilibrium constant for this system, called the solubility product expression, which relates the solubility product constant, Ksp, to the concentrations of the ions in the saturated solution;
+2 - 2 Ksp = [Pb ][I ] (eq. 2)
In more general terms, for the generic insoluble salt A2B3, whose aqueous equilibrium would be written as
+3 -2 A2B3(s) 2A (aq) + 3B (aq) (eq. 3) and the corresponding expression for Ksp would be written as
+32 -23 Ksp = [A ] [B ] (eq. 4)
Note that in each case, the solid is not included in the equilibrium constant expression, as the molar concentration of a pure solid or a pure liquid is considered to be a constant and is therefore part of the equilibrium constant for the system.
The reaction you will examine in this experiment is the solubility equilibrium involving calcium iodate, Ca(IO3)2, and the pertinent equations are thus
+2 - Ca(IO3)2 (s) Ca (aq) + 2 IO3 (aq) (eq. 5)
+2 - 2 Ksp = [Ca ][IO3 ] (eq. 6) 2
You will use titration data to determine the concentration of iodate ions in a saturated solution of calcium iodate, and from this information determine the value of Ksp for Ca(IO3)2. Once the iodate concentration is known, the calcium ion concentration is easily determined in this case since all the calcium ions came from dissolution of the calcium iodate. Thus the concentration of calcium ions is (stoichiometrically) half the concentration of the iodate ions. Once both ion concentrations are known, equation 6 will allow calculation of Ksp for this salt.
The Reactions
- The concentration of iodate ion, IO3 , will be determined by titration with a standardized solution of sodium thiosulfate, Na2S2O3, in the presence of potassium iodide, KI. Starch will be used as an indicator, and a very sharp blue-to-colorless transition of the titration mixture will mark the endpoint. The relevant reaction equations are summarized as follows.
− − + IO3 (aq) + 5I (aq) + 6H3O (aq) 3I2(aq) + 9H2O(l) (eq. 7)
This step, which occurs after adding both solid KI and aqueous acid to aliquots of saturated calcium iodate solutions, has the net effect of converting iodate ions to aqueous iodine, I2. Thiosulfate ion then reacts with aqueous iodine according to:
-2 − -2 I2(aq) + 2S2O3 (aq) 2I (aq) + S4O6 (aq) (eq. 8)
The overall titration reaction can be obtained by adding equations 7 and 8, then balancing in terms of both mass and charge: − − + IO3 (aq) + 5I (aq) + 6H3O (aq) 3I2(aq) + 9H2O(l) -2 − -2 3I2(aq) + 6S2O3 (aq) 6I (aq) + 3S4O6 (aq)
− -2 + − -2 IO3 (aq) + 6S2O3 (aq) + 6H3O (aq) I (aq) + 3S4O6 (aq) + 9H2O(l) (eq. 9)
Since the concentration of thiosulfate will be known and its volume consumed will be determined by titration, a simple stoichiometry calculation allows determination of the concentration of iodate in the original solution titrated.
Procedure
1. Obtain approximately 40 mL of the saturated calcium iodate solution in a small beaker. Rinse a clean buret with several small (~3-5 mL) portions of the solution with which the buret is to be filled, in this case the saturated calcium iodate. Be 3
sure to rinse the buret tip as well as the inside of the buret. Drain the buret completely, close the stopcock, and fill to near the top with the saturated calcium iodate solution. You do not have to start with the liquid level at 0.00 mL on the buret. Record on Data Sheet 1 the initial buret reading. All buret readings should be reported to two digits to the right of the decimal, i.e. ±0.01 mL.
2. Prepare another buret with standardized sodium thiosulfate solution in the same fashion as you did the calcium iodate buret. Record on Data Sheet 1 the exact concentration of this solution and the initial buret reading. It does not have to be at 0.00 mL.
SAFETY NOTE!!
Hydrochloric acid, HCl, is a strong acid. Avoid contact with your skin, eyes, and clothing. If you do spill some on yourself, wash immediately with cold water. Notify your instructor.
3. Carefully add a Teflon-coated magnetic stir bar to a clean 150 mL beaker. (See Figure 1.) Record the initial reading of the calcium iodate buret on Data Sheet 1. Deliver 10.00 mL of the saturated calcium iodate solution from the buret into the beaker. Record the final buret reading on Data Sheet 1. Add approximately 20 mL deionized water to the beaker, and swirl the beaker to mix the solutions. Add about 1 g. solid KI to the beaker, then 20 drops of 2 M HCl. Swirl to mix the contents, obtaining an orange-red homogeneous solution. Finally, add Figure 1 approximately 30-40 drops of 0.2% starch solution. The mixture should now be a dark greenish-black color.
4. Position the beaker on top of a magnetic stir plate such that the buret containing the sodium thiosulfate solution is directly above the beaker. Turn on the stirrer to obtain a vigorous mixing of the solution without splashing the contents up the inside of the beaker.
5. Record the initial buret reading of the sodium thiosulfate on Data Sheet 1. Begin titrating. As you titrate, the color of the solution in the beaker will change gradually to a dark blue-black. Titrate the contents of the beaker with the sodium thiosulfate solution until one drop or less causes the color to change from deep blue-black to colorless. Record the final buret reading on Data Sheet 1.
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6. Repeat steps 3 and 4 two more times with fresh samples of saturated calcium iodate solution. Record all data on Data Sheet 1. You should now have data for three trials.
HINT: If you use nearly the same volume (10.00 mL) of saturated calcium iodate for all trials, then the second and subsequent titrations can be done very quickly, because you will know approximately how much sodium thiosulfate will be required. For example, if the first trial required 21.69 mL thiosulfate, and the second and third trials use very close to the same volume of saturated calcium iodate as the first trial, then you can safely add very quickly at least 20.50 mL of sodium thiosulfate to the second and third trials, and then titrate very carefully from that point to a precise endpoint.
Calculations
1. Calculate the volumes of both the thiosulfate solution and the calcium iodate solution used, and record the results on Data Sheet 1. Convert both volumes to liters and record on Data Sheet 1 as well.
2. Using the molarity of sodium thiosulfate and the volume used for the titration, calculate the number of moles of thiosulfate ion reacted during the titration. Remember, moles of solute = molarity x volume in liters
Record your result on Data Sheet 1.
3. Calculate the number of moles of iodate which were present in the sample of saturated calcium iodate that you titrated. Refer to eq. 9 to obtain a mole ratio of -2 - S2O3 to IO3 to use in this stoichiometry calculation. Record your result on Data Sheet 1.
4. Calculate the molarity of iodate ions in the saturated calcium iodate solution. Remember that moles IO - molarity, M = 3 liters of sat'd Ca(IO32 ) solution
5. Calculate the molarity of calcium ions in the saturated solution of calcium iodate - that you titrated. Referring to eq. 5, there are two moles of iodate, IO3 , present for every one mole of calcium ions, Ca+2. Thus the molar concentration of calcium ions in the saturated solution is half the molarity of iodate ions. Record your result on Data Sheet 1.
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6. Record on Data Sheet 1 the molar solubility of calcium iodate. (Hint: which ion concentration will be equal to the molar solubility of the salt?)
7. Using the concentrations of both ions calculated above, now calculate Ksp for Ca(IO3)2 and record your result on Data Sheet 1.
8. Repeat calculations 1-7 for your other titrations.
9. Calculate the average value of Ksp from your data, and record the value on data Sheet 1.
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Solubility Product Constant of Calcium Iodate
Data Sheet 1
Trial #1 Trial #2 Trial #3
Molarity of Na2S2O3 (M) ______
Final Ca(IO3)2 buret reading (mL) ______
Initial Ca(IO3)2 buret reading (mL) ______
Volume of Ca(IO3)2 used (mL) ______
Volume of Ca(IO3)2 used (L) ______
Final Na2S2O3 buret reading (mL) ______
Initial Na2S2O3 buret reading (mL) ______
Volume Na2S2O3 used (mL) ______
Volume Na2S2O3 used (L) ______
Moles Na2S2O3 used ______
Moles iodate titrated ______
− Equilibrium molarity of IO3 (M) ______
Equilibrium molarity of Ca+2 (M) ______
Molar solubility of Ca(IO3)2 (M) ______
Ksp of Ca(IO3)2 ______
Average Ksp of Ca(IO3)2 ______7
Solubility Product Constant of Calcium Iodate
Post-Lab Questions
1. How would you expect the solubility of Ca(IO3)2 to be affected by the presence of some extra calcium ions. In other words, would Ca(IO3)2 be more soluble in 0.10 M Ca+2 solution or in pure water? Explain your answer.
2. Using the average value of your experimentally determined molar solubility, calculate the solubility of calcium iodate, Ca(IO3)2, in grams per liter of solution. Show all calculations.
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Solubility Product Constant of Calcium Iodate
Pre-laboratory Assignment
1. Complete the following equations, and write expressions for the solubility product constant, Ksp, for each of the following salts;
a) Cu2S(s)
b) Bi2S3(s)
c) Be(OH)2(s)
d) La(IO3)3(s)
e) Mn(OH)2(s)
+ -2 2. For the solubility equilibrium Ag2SO4(s) 2Ag (aq) + SO4 (aq), if the concentration of sulfate ion in a saturated solution of Ag2SO4 is found to be 0.0155 M, what is the value of Ksp for this salt?
-8 3. Calculate the molar solubility of PbF2 given that Ksp for this salt is 3.6 x 10 . (Hint: Write a balanced equilibrium equation first.)