Developing a Mastery of the Solubility Product Constant

Learning Objectives
• Be able to determine the solubility product constant for calcium iodate by titration of the iodate ion in a saturated solution.
• Be able to set up an equilibrium expression for a slightly soluble salt and write out the expression for K_sp in terms of concentrations of products and reactants.
• Given the solubility of a slightly soluble salt, in either mol/L or g/L, calculate the K_sp.
• Given the K_sp of a slightly soluble salt, calculate the molar solubility.
• Given a list of slightly soluble salts and their K_sp values, be able to list the salts in order of increasing solubility.
• Given a list of slightly soluble salts and their K_sp values, be able to determine the concentration of ions in solution.

Why Study Solubility?
Solubility data has many important technological applications in various fields of science. In designing drugs for medicinal purposes, a knowledge of solubility is required to facilitate drug delivery systems, as well as controlling the solubility of the drug in the blood stream. Engineers utilize theory to design water treatment facilities, which remove hazardous chemicals from drinking water. In addition, environmental chemists can explain pollutants by carefully analyzing solubility data. Our solubility analysis begins with analyzing titration data to determine molar solubility and the solubility product constant (K_sp) and will continue in subsequent labs with manipulating the solubility of a slightly soluble salt by analyzing the common ion effect, pH, complex ion formation, and amphoterism.

Procedure Hints

Step #2. Similar to Figure 2, assemble a ring, clay triangle, and funel with filter in order to perform a gravity filtration. Fold a piece of filter paper in half, then fold it in half again. In order to fit the filter paper snugly inside the funnel, tear off a small part of the corner as shown in the video below.

Open the filter paper so that one side of the funnel is in contact with a single layer of the filter paper and the other side of the funnel is in contact with three layers of the filter paper.It seems to work best if the side with the tear has three layers. The paper should be in smooth contact with the funnel at the top, but not near the point. Squirting in water with a wash bottle sometimes helps to keep the filter paper in place.

Step #12. To your TITRATION FLASK, add approximately 50 mL of water using a graduated cylinder and 2 g of solid KI.

In addition, add about 10 mL of 3 M HCl and swirl. This is converting all of the IO₃^–(aq) to I₂(aq), which will turn the solution brown.

This is consistent with the following equation:

IO₃⁻(aq) + 5 I⁻(aq) + 6 H⁺(aq) → 3 I₂(aq) + 3 H₂O(l).

Step #14. In each of your titrations, fill the 25-mL buret with 0.05XX M Na₂S₂O₃ and record your initial buret reading. Begin your titration and as the titration proceeds you will observe the brown color of the I₂ solution lighten to a more yellow color.

When you observe a yellow color stop the titration and prepare a fresh starch solution by adding 0.6 grams of vitex starch to roughly 30 mL of distilled water in a separate 250-mL Erlenmeyer flask.

Step #15. Add about 4 mL of the starch indicator solution to the TITRATION FLASK and in doing so, you should observe a blue (or dark brown) solution.

Step #16. Continue the titration until you reach the endpoint, which is the moment the solution turns colorless.

This end point indicates all of the iodine has been used up, which is consistent with the following equation:
3 I₂(aq) + 6 S₂O₃²⁻(aq) → 6 I⁻(aq) + 3 S₄O₆²⁻(aq)

By recording the final buret reading at the endpoint you are able to calculate the number of moles of S₂O₃²⁻(aq) required in the titration.

At the endpoint of a titration, chemists can convert the moles of a substance that is known to the moles of something that is unknown. Most titrations performed up to this point involved neutralization reactions between acids and bases. At the end point of an acid/base titration, the moles of acid are equal to the moles of base. The chemical reaction in this experiment does not involve a neutralization reaction between an acid and a base. What we are observing is the following overall redox reaction:

IO₃⁻(aq) + 6 H⁺(aq) + 6 S₂O₃²⁻(aq) → I⁻(aq) + 3 S₄O₆²⁻(aq) + 3 H₂O(l)

From this overall chemical equation, we can come up with the most important relationship in this particular experiment. At the endpoint of the titration:

6 mol S₂O₃²⁻(aq) = 1 mol IO₃⁻(aq)

This will allow you to calculate the molarity of IO₃⁻ for your titrations.