Materials Included In Kit

Additional Materials Required

Prelab Preparation

Saturated solution must be freshly prepared. Add 150 mL of distilled or deionized water directly to the sample bottle of calcium iodate. Stir the resulting solution with a glass stirring rod. Allow the mixture to stand for 30 minutes with occasional stirring. Decant the liquid from the solid through a piece of folded filter paper and store the saturated calcium iodate solution in a stoppered bottle or flask. Do NOT introduce any extraneous water to the saturated solution during filtration, storage or transfer of the solution.

Safety Precautions

Hydrochloric acid solution is toxic and corrosive to eyes and skin tissue. Calcium iodate is a mild oxidizer and a body tissue irritant. Sodium thiosulfate is slightly toxic by ingestion and a body tissue irritant. Avoid contact of all chemicals with eyes and skin. Wear chemical splash goggles, chemical-resistant gloves and a chemical-resistant apron. Please review current Safety Data Sheets for additional safety, handling and disposal information.

Disposal

Please consult your current Flinn Scientific Catalog/Reference Manual for general guidelines and specific procedures, and review all federal, state and local regulations that may apply, before proceeding. All of the solutions can be disposed of according to Flinn Suggested Disposal Procedure #26b.

Teacher Tips

• Enough materials are provided in this kit for 30 students working in pairs or for 15 groups of students. The laboratory procedure can reasonably be completed in one 50-minute class period. The Preab Activity may be assigned separately as preparation for lab, or it can be used as the basis of a cooperative class discussion.
• Preparation of the saturated calcium iodate solution requires about 45 minutes of teacher prep time prior to class. This may be done 1–2 days before lab if the resulting solution is stored in a stoppered bottle or flask.
• Remind the students that the sample being analyzed is a saturated solution and that they must carefully avoid adding any extra water to it prior to measuring and transferring a known sample volume to their titration beaker.
• It is not unusual to encounter large errors in solubility product experiments. The most important sources of error in solubility versus K_sp measurements are ones of omission—neglecting to take into account partial dissociation of ionic compounds, ion-pair formation, hydrolysis of anions, complex-ion formation involving the metal ions, etc. Almost all of these factors increase the solubility of compounds in solution compared to their K_sp calculations.
• Several Flinn student laboratory kits have been developed to allow students to further probe their understanding of the solubility of ionic compounds. Precipitation Reactions and Solubility Rules (AP4862) allows students to analyze solubility patterns and formulate the rules of solubility. Factors Affecting Solubility (AP4865) provides an inquiry-based activity to investigate the factors that affect the rate at which substances dissolve in solution.

Answers to Prelab Questions

1. The ionic compound magnesium hydroxide, which is found in the over-the-counter consumer product Milk of Magnesia™, has the formula Mg(OH)₂. Magnesium hydroxide dissolves in water according to the following equation:
   {11941_PreLab_Equation_1}

   Write the solubility product expression for magnesium hydroxide.

   K_sp = [Mg²⁺][OH^–]²
2. The K_sp value for magnesium hydroxide is 5.61 x 10^–12 at 25 °C. In a saturated solution of magnesium hydroxide, the magnesium ion concentration was found to be 1.1 x 10^_–4 M. Calculate the hydroxide ion concentration in the solution.

   (1.1 x 10^–4)[OH^–]² = 5.61 x 10^–12
   [OH^–]2 = (5.61 × 10–12)/(1.1 × 10–4)
   [OH–] = 2.26 × 10–4 M

Sample Data

Answers to Questions

1. Determine the mean value of the solubility product for calcium iodate and the average deviation for your three trials.
   
   {11941_Answers_Equation_1}
2. The literature value of the Ksp for calcium iodate is 6.47 x 10^–6 at 25 °C. Use the following equation to calculate the percent error in your determination of the solubility product for calcium iodate. Note: Do not be alarmed if your average deviation is very large. This is not uncommon in K_sp experiments, even under the most careful conditions.
   
   {11941_Answers_Equation_2}
   {11941_Answers_Equation_3}
3. A student did not notice that the test tube in which she obtained her saturated calcium iodate solution had some water in it. Discuss the effect of this error on the calculated value of the solubility product—was her K_sp value likely to be higher or lower as a result?
   The presence of extraneous water in the test tube will cause dilution of the saturated solution of calcium iodate. This will result in a lower value of the measured concentration of iodate ion in solution, and by extension, a lower than expected value for the calculated K_sp of calcium iodate.
4. Calcium oxalate is a sparingly soluble ionic compound composed of calcium ions (Ca^₂₊) and oxalate ions (C₂O₄^2–). It is a main culprit in kidney stone disease, due to the formation of calcium oxalate crystals in the urine. A typical value for the concentration of calcium in urine is 0.3 grams per liter. Given a K_sp value for calcium oxalate of 1.3 x 10^–8, calculate the minimum concentration of oxalate ion in urine that could lead to precipitation of calcium oxalate.
   
   {11941_Answers_Equation_4}

References

Special thanks to Bob Lewis, Downers Grove North High School in Downers Grove, IL, and John Little, St. Mary’s High School in Stockton, CA, for the lab idea and sample procedure.

Introduction

Gradually, over thousands of years, a cave is carved out of limestone rock. How does this happen? The mineral calcium carbonate first dissolves and then crystallizes out again in the form of stalactites and stalagmites. What are the factors that govern the solubility of an ionic compound in water? Is it possible to measure how much solid will dissolve?

Concepts

• Solubility equilibrium
• Solubility product constant (K_sp)

Background

When an ionic compound is placed in water, an equilibrium occurs between the undissolved solid compound and dissolved aqueous cations and anions. The solubility equilibrium that results is described quantitatively in terms of the equilibrium constant for this reversible reaction. Since the reaction involves solubility, the equilibrium constant is called the solubility product constant (K_sp).

Consider the following equilibrium that results when calcium iodate, a sparingly soluble ionic compound, dissolves in water (Equation 1). The solubility product expression for this equilibrium reaction is shown in Equation 2.

The purpose of this lab is to determine the solubility product constant of calcium iodate at room temperature. The concentration of iodate ion in a saturated solution of pure calcium iodate will be measured experimentally by titration with sodium thiosulfate. The calcium ion concentration in the saturated solution is related to the iodate concentration by the formula of the ionic compound. Substituting these two concentration values into Equation 2 makes it possible to calculate the K_sp value for calcium iodate.

The titration reaction that takes place is actually a combination of two reactions. In the first reaction, excess potassium iodide is added to a saturated solution of calcium iodate to reduce all of the iodate ion to triiodide, I₃^– (Equation 3). In the second reaction, sodium thiosulfate is added to convert the triiodide to iodide ion (Equation 4). Equation 6 gives the overall relationship between the number of moles of sodium thiosulfate added and the number of moles of iodate ion present in the original solution. Equation 6 was obtained by multiplying Equation 4 by a factor of three (to give Equation 5) and then adding Equations 3 and 5, as shown below. (Hint: Equation 4 is multiplied by three so that the number of moles of I₃^– consumed by Equation 5 equals the number of moles of I₃^– produced by Equation 3.) Sodium thiosulfate solution of known concentration (in moles of sodium thiosulfate per gram of solution) is prepared. The mass of the solution is measured and is added dropwise to the iodate/iodide reaction mixture until all of the triiodide ion has been consumed and no further reaction is observed. This point is called the endpoint of the titration. A starch indicator is added to the solution to provide a visible color change that can be taken as a sign that the endpoint has been reached. Starch forms deep blue- or purple-colored complexes with I₃^–. Disappearance of the purple color is very abrupt and signals that the endpoint has been reached.

At the endpoint of the reaction the number of moles of thiosulfate added is related to the number of moles iodate ion present initially by the stoichiometry of Equation 6. The amount of sodium thiosulfate solution that is required to reach the endpoint is measured by obtaining the mass of the solution bottle before and after the titration reaction. The difference in mass corresponds to the mass of sodium thiosulfate solution added. Since the concentration of the sodium thiosulfate solution is also known—based on how it was prepared—the number of moles of thiosulfate can be calculated and related to the concentration of iodate ion in a saturated solution of calcium iodate.

Materials

Prelab Questions

1. The ionic compound magnesium hydroxide, which is found in the over-the-counter consumer product Milk of Magnesia™, has the formula Mg(OH)₂. Magnesium hydroxide dissolves in water according to the following equation:
   {11941_PreLab_Equation_1}
   Write the solubility product expression for magnesium hydroxide.
   
   
2. The K_sp value for magnesium hydroxide is 5.61 x 10^–12 at 25 °C. In a saturated solution of magnesium hydroxide, the magnesium ion concentration was found to be 1.1 x 10^–4. Calculate the hydroxide ion concentration in the solution.

Safety Precautions

Hydrochloric acid solution is toxic and corrosive to eyes and skin tissue. Calcium iodate is a mild oxidizer and a body tissue irritant. Sodium thiosulfate is slightly toxic by ingestion and a body tissue irritant. Avoid contact of all chemicals with eyes and skin. Wear chemical splash goggles, chemical-resistant gloves and a chemical-resistant apron. Wash hands thoroughly with soap and water before leaving the laboratory.

Procedure

Preparation

1. Mass out 0.20–0.22 grams of sodium thiosulfate pentahydrate in a tared 50-mL beaker. Record the exact mass on the Data Sheet.
2. Add 10 mL of distilled water and record the mass of the solution on the Data Sheet.
3. Stir until fully dissolved and then transfer the solution to a 15-mL dropping bottle. This is called the weight buret.

Titration Analysis

4. Record the initial mass of the weight buret (dropping bottle and the solution, see step 3) in the Titration Table.
5. Add about 20 mL of water to a 50-mL beaker. Use a 5-mL serological pipet to add about 2.0 mL of saturated calcium iodate solution to the beaker. Record the exact volume in the Titration Table.
6. Add 10 drops of 2 M HCl to the solution in the beaker, followed by 0.2 g of potassium iodide. Stir until dissolved. The solution should be brown due to the formation of triiodide ion.
7. Remove about 1 mL of the solution using a Beral pipet. Set the Beral pipet aside for later use (store it upside down in a secure place). This pipet is called the titration thief.
8. Add sodium thiosulfate solution dropwise from the weight buret to the calcium iodate solution in the beaker until a light yellow color is reached.
9. Add 10 drops of starch solution to the beaker. The solution should be deep purple.
10. Continue adding sodium thiosulfate dropwise from the weight buret to the solution in the beaker until the purple solution just turns colorless. Swirl the solution in the beaker carefully to mix all of the reactants.
11. Transfer the solution from the titration thief back to the beaker.
12. Continue adding thiosulfate solution dropwise very carefully from the weight buret until the solution in the beaker is again colorless. Note: Ideally, one drop of thiosulfate solution should change the reaction mixture to colorless.
13. Mass the weight buret and record its final mass in the Titration Table.
14. Repeat steps 5–13 for two additional trials.
15. Consult your teacher for proper disposal information.

Student Worksheet PDF

11941_Student1.pdf