Solubility of Lead (II) Iodide

To determine the solubility product constant, Ksp, for PbI2.
To create and utilize a photometric titration curve.

[This laboratory exercise was adapted from an article by Gary W. Rice "Potentiometric and Photometric Methods for Determining the Solubility of Lead Iodide" in the Journal of Chemical Education, May 1990]

Lead (II) iodide is a slightly soluble salt. Its solubility can be determined by measuring either the lead ion concentration or the iodide ion concentration in a saturated solution. In this lab we will use a redox titration to measure the iodide ion concentration. This investigation is more involved than past investigations because not one but two redox reactions will be occurring each with its own endpoint. Data will be collected and a photometric titration curve of absorbance vs mL of reactant plotted in order to determine the endpoint of the titration. The redox reactions are:

As a saturated solution of lead (II) iodide is titrated with Ce4+ [as cerium (IV) nitrate] the concentration of free iodine (I2) will first rise and then begin to fall. The change in concentration of the I2 can be monitored by measuring the change in absorbance as the concentration of the iodine changes. Absorbance changes will be monitored at a wavelength of 435 nM. Ideally a plot of Absorbance vs mL of Ce4+ added should look like the following:

In theory either endpoint can be used in determining the concentration of iodide ion present. However past experience in this class has indicated that only the first endpoint is reliable.

1. Prepare solid lead iodide by mixing together about 10 mL of 0.1M lead (II) nitrate and 20 mL of 0.1M potassium iodide. Centrifuge the precipitate and wash several times to insure that it is free of all electrolytes. Let the solid lead iodide stand in contact with about 150 mL of water for a week before proceeding. The solution should be shaken every day. (Alternatively, you can heat the solution to about 80-85 degrees Celsius to dissolve most of the lead (II) iodide. Then let the solution stand for 24 hours so that the excess lead (II) iodide precipitates out. When filtered the solution should now be saturated in lead (II) iodide at room temperature.)

2. Extract 50.0 mL of the saturated lead (II) iodide, and place in a beaker. Be careful that no solid lead (II) iodide is transferred. Filter if necessary. To the lead (II) iodide solution add 50.0 mL of 2 N hydrochloric acid. Titrate the saturated lead (II) iodide solution with 0.05 M Ce4+ by adding 0.50 mL portions. After adding each portion of the titrant transfer enough of the solution to a cuvette and measure the absorbance at 435 nM. Return the iodide solution to the beaker before adding the next portion of titrant.

1. Make a plot of absorbance vs. mL of titrant added and determine the two endpoints.

2. Determine the mmols of titrant added and the mmols of iodide ion present in the saturated lead (II) iodide solution.

3. Calculate the Ksp of lead (II) iodide.

4. How would the calculated value of Ksp be affected if a small solid piece of lead (II) iodide were accidentally transferred to the beaker before the titration was done? Explain.

5. Ideally each absorbance reading should have been corrected for the dilution that took place because of the titrant. Extra Credit: Make these corrections and recalculate the Ksp . In your opinion does the dilution have much effect on the calculated Ksp ?

6. From the class values listed on the chalkboard for the Ksp of lead (II) iodide determine which ones would be considered outliers at the 96% confidence level and rejected. After rejecting these values calculate the class average for the and the standard deviation.

Teacher Notes

Enhanced for Netscape 2.0.

This is nice laboratory exercise and gives students a chance to work with a spectrophotometric endpoint. The cerium (IV) nitrate solution must be fresh.

Questions? Comments??