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Chemistry 212 Lab

Chemistry 212 Lab Fall 2004 

Solubility Product Constant (Ksp) and the Common-Ion Effect

for a Salt of Limited Solubility


      Determine the solubility product constant (Ksp) for a sparingly soluble salt.  Study the effect on the aqueous solubility equilibrium of this salt in the presence of an additional amount of one of the ions that it contains (in this case, the effect of added Ca2+). 


      This experiment deals with saturated aqueous solutions of salts with limited solubility in water.  Under ordinary conditions, it might be fair to call the material you'll use today as insoluble, or at best slightly soluble.  At such conditions of limited aqueous solubility, the solid salt and its ions in solution are in a state of dynamic equilibrium between the solid phase and the solution phase.  Concentrations of ions in solution are exceedingly low, and care must be taken to obtain good results, when performing the experimental procedure outlined shortly.

      As an example of a slightly soluble salt, consider a saturated solution of silver sulfate, Ag2SO4.  In this example, the equilibrium lies far to the left, owing to the fact that silver sulfate is only slightly soluble in water:

Ag2SO4(s) <-=====-> 2Ag+(aq)  +  SO42(aq

The corresponding equilibrium constant expression for this system is: 

                                        [Ag+]2 [SO42]

                              K  =  --------------------


Since the solid component has a near-constant concentration, it is not used in the equilibrium expression, simplifying it to the solubility product

                K = [Ag+]2 [SO42

The new symbol Ksp is called the solubility product constant.  Most solubility product constants have very small values, as is the case for silver sulfate, where Ksp = 1.1 x 105

                1.1 x 105 = [Ag+]2 [SO42

In 1 liter of a saturated, aqueous solution of silver sulfate, the molar concentration of silver ion can be found by letting x equal the sulfate ion concentration, 2x the silver ion concentration, and substituting into the solubility product equation the known value for Ksp

1.1 x 105 = [Ag+]2 [SO42

    1. x 105 = [2x]2 [x]

1.1 x 105 = 4 x3 

2.75 x 106 = x3 

x = [SO42] = 1.4 x 102 M 

2x = [Ag] = 2.8 x 102 M


      The equilibrium process in this experiment is a saturated aqueous solution of calcium iodate, Ca(IO3)2.  The relevant solubility equation and solubility product expression, are both shown below. 

Ca(IO3)2(s) <-=====-> Ca2+(aq)  +  2IO3(aq

                Ksp = [Ca2+] [IO3]2 

For a saturated solution of calcium iodate, if you can determine either the molar concentration of calcium ion, or the molar concentration iodate ion, the solubility product constant can be found using the reverse of the process shown above. There was found the silver ion concentration, in a saturated aqueous solution, from a known value for Ksp.  In other words, if the calcium ion concentration in today's experiment was found to be 0.1 M, you could immediately say the concentration of iodate ion must be half that value, or 0.05 M, according to the stoichiometry of the solubility equation given above.  The solubility product constant could then be found with simple arithmetic.  In this experiment, the iodate ion concentration of a saturated calcium iodate solution will be found via a redox titration with sodium thiosulfate, Na2S2O3.

      Also, the effect of adding a source of additional calcium ion on the solubility of calcium iodate will be determined.  This is accomplished via redox titration of a second calcium iodate solution containing an additional source of aqueous calcium ion (e.g., added CaCl2).  It is possible to predict the result qualitatively upon the calcium iodate equilibrium, before actually performing the experiment.

      According to LeChâtelier's principle, adding an additional source of calcium ion (besides that of calcium iodate itself) should cause the solubility equilibrium to shift to the left, favoring more solid calcium iodate.  This is an example of the common-ion effect.  The common-ion effect, in this experiment, should lead to a reduced solubility of calcium iodate, and a corresponding change in the solubility product constant.


Details of the Redox Titration of Iodate Ion 

      The concentration of iodate ion (IO3-) will be determined by titration with a standardized sodium thiosulfate (Na2S2O3) solution in the presence of potassium iodide (KI).  Starch will be used as an indicator, and a sharp blue-to-clear transition will mark the equivalence point.  The relevant reaction equations are summarized as follows. 

IO3(aq)  +  5I(aq)  +  6H3O(aq)  -------->  3I2(aq)  +  9H2O(l

This step, which occurs after adding both solid KI, and aqueous acid, to aliquots of saturated iodate solutions, has the net effect of converting iodate ions to aqueous iodine.  Thiosulfate ion then reacts with aqueous iodine according to:

I2(aq)  +  2S2O32(aq)  -------->  2I(aq)  +  S4O62(aq

The net titration reaction can be obtained by combining the two reactions above, then balancing for mass and charge:

IO3(aq)  +  5I(aq)  +  6H3O(aq)  -------->  3I2(aq)  +  9H2O(l

3I2(aq)  +  6S2O32(aq)  -------->  6I(aq)  +  3S4O62(aq)


IO3(aq)  +  6S2O32(aq)  +  6H3O(aq)  -------->  I(aq)  +  3S4O62(aq)  +  9H2O(l) 

It is important to notice the molar ratio of iodate ion to thiosulfate ion in the above complete, balanced titration reaction.  Ask your instructor if the thiosulfate solution available for you to use has been standardized, or whether you will have to standardize the thiosulfate solution as part of the experimental procedure.  Aliquots from prepared, saturated solutions of calcium iodate in water, with and without additional calcium ion, will be titrated according to the reaction chemistry of the equations outlined above. 

Experimental Notes

•  It is important to handle the saturated iodate solutions with great care, so as not to allow any of the solid calcium iodate into you titrations.  This will obviously lead to a higher value for concentration of iodate than is actually present in the saturated solutions.

•  It is recommended that you perform 4 titrations of unknown calcium iodate after standardizing the thiosulfate solution.  The first of these should be a 'scout', where 1 mL portions of thiosulfate solution are added quickly to determine the approximate volume needed to reach equivalence.

•  You may find it convenient to set up all four calcium iodate unknown determinations at once.  This is fine to do, as the titration solutions, once prepared, should remain stable until you are ready for them.

•  Titrated solutions, and any left over thiosulfate or iodate solutions, are safe to dilute and rinse down the sinks.


Experimental Procedure 

Note: it is important for the Erlenmeyer flasks and other glassware you'll use to be as clean as possible when performing these redox titrations.  Please take a moment to at least give them a thorough rinsing--they need not be completely dry before proceeding. 

Saturated solutions of calcium iodate in water, and calcium iodate with added calcium ion, are available in the hoods.  Please take some care to not disturb the solid calcium iodate present in the bottles when obtained your portions of solutions to use. 

Obtain approximately 40 mL of the two different calcium iodate solutions in small beakers.  One is labeled "Calcium Iodate, Aqueous" and the other "Calcium Iodate, Added Calcium Ion".  Set up a 50 mL buret and stand, fill the buret with standardized sodium thiosulfate solution, and take a reading of initial volume to two decimal places.  Be sure and record the exact concentration of the thiosulfate solution that is listed on the reagent bottle.  This solution has been standardized by the Preparation Lab, and its concentration should be printed on the label.  You will eventually use about 150 mL of this standardized thiosulfate solution. 

General Procedure for Iodate Solution Preparation 

Pipet 10.0 mL of a calcium iodate saturated solution into a clean 125 mL Erlenmeyer flask.  Add roughly 20 mL of distilled water to this saturated iodate solution, swirl to mix.  Dissolve ~1 g of solid KI into the iodate/water solution, then add 20 drops of 2 M HCl from a dropper bottle.  Swirl to mix the contents, obtaining dark red homogeneous solutions.  To complete the sample preparation, add approximately 30-40 drops of a 0.2% starch solution.  The titration solutions should be a dark blue-black color. 

You will need to prepare 2 titration solutions of the "Calcium Iodate, Aqueous", and 3 titration solutions of the "Calcium Iodate, Added Calcium Ion", for a total of 5 samples.  It is not necessary to always use a 125 mL Erlenmeyer flask--anything that will hold this volume, provided it is clean, can be used. 

Performing the Titrations 

The first titration is a "scout" using an "Added Calcium Ion" sample, intended to make the remaining 4 titrations go faster and with greater accuracy.  Perform the scout titration by adding 1 mL of the thiosulfate solution at a time, swirling and noting any changes in color.  Remember that the starch/iodine endpoint is a sharp blue-black to clear transition, and it is unlikely you will miss it.  Continue adding 1 mL portions of thiosulfate solution, until 1 mL added causes the solution to become clear and colorless.  Assuming that all the titration solutions to follow have been prepared in the same manner as the scout, this gives an idea of what the actual equivalence volume is to +/ 1 mL. 

For the remaining titrations, it is permissible to add enough thiosulfate solution so that you are within 2 mL of the equivalence volume.  Be sure and record initial volumes before titrating, and final volumes at equivalence, from the barrel of the buret to 2 decimal places. 


Using the above scout procedure, imagine the equivalence volume was determined to be after adding 12 mL of thiosulfate solution; but before adding 13 mL.  This means that 12 1 mL additions failed to produce a color change, but the 13th turned the solution from blue to clear. 

For the next titration, an accurate initial volume of 0.06 mL was read, the titration was performed by adding 10 mL of the thiosulfate solution all at once to begin.  If the samples were prepared identically, this should be before enough thiosulfate has been added to reach the equivalence volume.  From this point, the titration continues on a drop by drop basis.  Eventually, one single drop was added that caused the color of the titration solution to change from blue to colorless.  The final volume reading was taken at that time to be 12.74 mL. 

The volume added to reach equivalence is therefore carefully determined, and found to be (12.74 mL  0.06 mL) = 12.68 mL.  This volume, along with the molarity of the thiosulfate solution, allows the most accurate possible determination of moles thiosulfate

It is a few steps from there to convert moles thiosulfate, to moles iodate ion, to concentration of iodate ion, finally to solubility product constant (Ksp) for calcium iodate.


Data Tables -- Solubility Product Constant/Common-Ion Effect 

Part A: Calcium Iodate, No Added Calcium Ion

                                                  Trial #1    Trial #2 

Concentration of Sodium Thiosulfate Solution (M)  __________  __________ 

Volume of Calcium Iodate Solution Added (L)  __________  __________ 

Final Volume, Thiosulfate Solution (mL)   __________  __________ 

Initial Volume, Thiosulfate Solution (mL)   __________  __________ 

Total Volume, Thiosulfate Solution (mL)   __________  __________ 

Total Volume, Thiosulfate Solution (L)   __________  __________ 

Moles Sodium Thiosulfate Used    __________  __________ 

Moles Iodate       __________  __________ 

Equilibrium Concentration of Iodate Ion (M)   __________  __________ 

Equilibrium Concentration of Calcium Ion (M)  __________  __________ 

Molar Solubility of Calcium Iodate (M)   __________  __________ 

Ksp, Calcium Iodate      __________  __________ 

Average For Ksp, Calcium Iodate           __________


Data Tables -- Solubility Product Constant/Common-Ion Effect 

Part B: Calcium Iodate, Added Calcium Ion

                                                  Trial #1    Trial #2 

Concentration of Sodium Thiosulfate Solution (M)  __________  __________ 

Volume of Calcium Iodate Solution Added (L)  __________  __________ 

Final Volume, Thiosulfate Solution (mL)   __________  __________ 

Initial Volume, Thiosulfate Solution (mL)   __________  __________ 

Total Volume, Thiosulfate Solution (mL)   __________  __________ 

Total Volume, Thiosulfate Solution (L)   __________  __________ 

Moles Sodium Thiosulfate     __________  __________ 

Moles Iodate       __________  __________ 

Equilibrium Concentration of Iodate Ion (M)   __________  __________ 

Equilibrium Concentration of Calcium Ion (M)  __________  __________ 

Molar Solubility of Calcium Iodate    __________  __________

(added Ca2 ion source)

Average Molar Solubility of Calcium Iodate          __________

(added Ca2 ion source) 

"Scout Titration" 

Volume of Calcium Iodate Solution Added (L)  __________ 

Final Volume, Thiosulfate Solution (nearest mL)  __________ 

Initial Volume, Thiosulfate Solution (nearest mL)  __________ 

Total Volume, Thiosulfate Solution (nearest mL)  __________ 

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