Materials Included In Kit

Additional Materials Required

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. Reaction solutions may be disposed of according to Flinn Suggested Disposal Method #26b. Unused sodium hydroxide solution may be disposed of according to Flinn Suggested Disposal Method #10. The unused calcium nitrate solution may be disposed of according to Flinn Suggested Disposal Method #26b.

Lab Hints

• Enough materials are provided in this kit for 30 students working in pairs, or for 15 student groups. Students will need approximately 45 minutes to complete the experiment.
• Students must use good technique to measure small solution volumes and make dilutions using Beral capillary pipets with microtips. Calculations involve determination of the concentrations of the diluted solutions and calculation of the solubility product.
• It is important to use good technique in delivering and counting the drops. Students need to be patient and repeat the process if an error occurs. Hold pipets vertically and be sure that there are no air bubbles in the barrel of the pipet.
• Equilibrium constants are sometimes recorded using units, and at other times the units are omitted. Choose the method used in the current textbook. However, students should recognize what the correct units would be if they were included.
• Be sure to review the calculations for determining the Ca²⁺ and OH^– concentrations in each of the wells.

Answers to Prelab Questions

1. Write a balanced chemical equation for calcium hydroxide, Ca(OH)₂, dissolving in water, and the solubility product expression for this equilibrium.
   
   {12980_Answers_Equation_6}

   K_sp = [Ca²+][OH^–]²
   

2. Explain how you determined the concentration of each ion in the first well where no precipitation appeared. Why is this the well that you use to find the solubility product?
   
   Solutions were initially 0.10 M. Dilution changes the concentration to 0.10 × (½)^n–1 M where n = the number of the well. Solution concentrations were halved one more time when the Ca²⁺ and OH^– solutions were mixed. It is assumed that the first well, where no precipitate appears, is a saturated solution.
   
   
3. For the calcium ion serial dilutions, calculate the concentrations of calcium ions, [Ca²⁺], before and after the addition of hydroxide ions, (OH^–) to each well.
   
   Before:
   [Ca²⁺] = 0.10 M × (½)^n–1, where n = well number
   Well 1 – [Ca²⁺] = 0.10 M × (½)^1–1 = 0.10 M
   Well 2 – [Ca²⁺] = 0.10 M × (½)^2–1 = 0.05 M
   Well 3 – [Ca²⁺] = 0.10 M × (½)^3–1 = 0.025 M
   Well 4 – [Ca²⁺] = 0.10 M × (½)³ = 0.013 M
   Well 5 – [Ca²⁺] = 0.10 M × (½)⁴ = 0.0063 M
   Well 6 – [Ca²⁺] = 0.10 M × (½)⁵ = 3.1 × 10^–3 M
   Well 7 – [Ca²⁺] = 0.10 M × (½)⁶ = 1.6 × 10^–3 M
   Well 8 – [Ca²⁺] = 0.10 M × (½)⁷ = 7.8 × 10^–4 M
   Well 9 – [Ca²⁺] = 0.10 M × (½)⁸ = 3.9 × 10^–4 M
   Well 10 – [Ca²⁺] = 0.10 M × (½)⁹ = 2.0 × 10^–4 M
   Well 11 – [Ca²⁺] = 0.10 M × (½)¹⁰ = 9.8 × 10^–5 M
   Well 12 – [Ca²⁺] = 0.10 M × (½)¹¹ = 4.9 × 10^–5 M
   
   After 5 mL of 0.10 M hydroxide ion, (OH^–), is added,
   All [Ca²⁺] values in each well are cut in half. The [OH^–] in each well is 0.1 M x ½ = 0.05 M
   
   
4. Repeat the calculations for the hydroxide ion and the calcium ion concentrations in the 12 wells with the hydroxide ion serial dilution. Record these final concentrations under OH^– dilutions.
   
   The concentration values are the same as in the calcium ion serial dilutions, with the values for [OH^–] and [Ca²⁺] reversed.

Sample Data

Calcium Ion Serial Dilutions

First well with no precipitation: Well 6

Concentration of Ca²⁺: 1.6 x 10^–3 mol/L

Concentration of OH^–: 5.0 x 10^–2 mol/L

K_sp[Ca(OH)₂]= 4.0 x 10^–6 M³

Hydroxide Ion Serial Dilutions

First well with no precipitation: Well 4

Concentration of Ca²⁺: 5.0 x 10^–2 mol/L

Concentration of OH^–: 6.3 x 10^–3 mol/L

K_sp[Ca(OH)₂]= 2.0 x 10^–6 M³

Answers to Questions

1. Calculate the concentration of Ca²⁺ ions and OH^– ions in the first well of the calcium ion dilution series with no precipitate. Using these concentrations, determine the solubility product, the K_sp of calcium hydroxide. Enter these values in the Data Table under Calcium Ion Serial Dilutions.
   
   Dilution of Ca(NO₃)₂: First well with no precipitate is well 6.
   Well 6: [Ca²⁺] = 0.10 × (½)⁵ M = 3.1 × 10^–3 M

   After NaOH is added: [Ca²⁺] = 1.6 × 10^–3 M
   [OH^–] after solutions are mixed = 0.050 M
   K_sp = [Ca²⁺] [OH^–]² = 1.6 × 10^–3 M × (0.050 M )² = 4.0 × 10^–6 M³

2. Calculate the concentration of calcium and hydroxide ions in the first well of the hydroxide ion dilution series where there is no precipitate, and again calculate the value of K_sp. Enter these values in the Data Table under Hydroxide Ion Serial Dilutions.
   Dilution of NaOH: First well with no precipitate is well 4.
   Well 4: [OH^–] = 0.10 × (½)³ M = 0.013 M

   After Ca(NO₃)₂ is added: [OH^–] = 6.3 × 10^–3 M
   [Ca²⁺] after solutions are mixed = 0.050 M
   K_sp = [Ca²⁺][OH^–]² = 0.050 M × (6.3 × 10^–3 M)² = 2.0 × 10^–6 M³

3. How did the values obtained from the two trials compare with each other?
   Values from the two trials compared within the limits obtainable with this technique. 
   
   
4. Does this method give values that are too low or too high? Why?
   Values obtained will probably be too low. The correct concentration is only known to be somewhere between the first well with no precipitate and the last well containing precipitate.
   
   
5. What would make the method more accurate?
   It would be better if more values were obtained using the concentrations between the last well containing precipitate and the first well containing no precipitate.
   
   
6. Would the results be better if the concentrations of the last well where precipitation occurred were averaged with the first well where there was no precipitate? Is there any justification for doing this? Try it!
   There is justification for this technique because the correct value is only known to be between the concentrations of the last well containing precipitate and the first well containing no precipitate.
   Using data from the dilution of calcium ions:
   

   Last well with precipitate is well 5.

   [Ca²⁺] in well 5 = 6.3 × 10^–3 M
   [Ca²⁺] in well 6 = 3.1 × 10^–3 M

   Average [Ca²⁺] = 4.7 × 10^–3 M. After dilution with OH^–, [Ca²⁺] = 2.4 × 10^–3 M
   K_sp = [Ca²⁺] [OH^–]² = 2.4 × 10^–3 M × (0.050 M)² = 6.0 × 10^–6 M³

   Using data from the dilution of hydroxide ions:
   

   Last well with precipitate is well 3.

   [OH^–] in well 3 = 0.025 M
   [OH^–] in well 4 = 0.012 M

   Average [OH^–] = 0.019 M. After dilution with Ca²⁺, [OH^–] = 9.4 × 10^–3 M
   K_sp = [Ca²⁺][OH^–]² = 0.050 M × (9.4 × 10^–3 M)² = 4.4 × 10^–6 M³

   Average K_sp = 5.2 × 10^–6 M³
   These values are closer to the accepted value of 5.5 × 10^–6 M³.

Introduction

This experiment uses microscale techniques to determine the solubility product, K_sp, of calcium hydroxide.

Concepts

• Solubility product
• K_sp
• Saturated solution
• Precipitation reaction

Background

The solubility product constant, K_sp, is an equilibrium constant. The equilibrium is formed when an ionic solid dissolves in water to form a saturated solution. The equilibrium exists between the dissolved ions and the undissolved solid. A saturated solution contains the maximum concentration of ions of the substance that can dissolve at the temperature of the solution.

For any solid in equilibrium with its ions in solution the general chemical equation is

where square brackets refer to molar concentrations of the ions. A knowledge of the K_sp of a salt is useful, since it is used to determine the concentration of ions of the compound in a saturated solution. Understanding the maximum concentration is important to determine when a compound will precipitate out of solution and when it will stay in solution.

The solubility product to be determined in this experiment is that of calcium hydroxide, Ca(OH)₂. Calcium hydroxide is a strong base, but is not very soluble in water.

Experiment Overview

This experiment uses a microscale technique to determine the solubility product of calcium hydroxide, Ca(OH)₂. Serial dilutions of the solution of calcium nitrate are placed in separate wells of a reaction strip. The diluted solutions are combined with an equal volume of the sodium hydroxide solution. Calcium hydroxide precipitates in the wells where the concentrations of calcium and hydroxide ions exceed the solubility product. The first well where no precipitate is present is assumed to be a saturated solution, and the ion concentrations are used to calculate the solubility product. The process is repeated using serial dilutions of the sodium hydroxide solution and equal volumes of the 0.10 M calcium nitrate solution.

Materials

Prelab Questions

1. Write a balanced chemical equation for calcium hydroxide, Ca(OH)₂, dissolving in water, and the solubility product expression for this equilibrium.
2. Explain how to determine the concentration of each ion in the first well where no precipitation appeared. Why is this the well that is used to find the solubility product?
3. For the calcium ion serial dilutions, calculate the concentration of calcium ions, [Ca²⁺], before and after the addition of hydroxide ions, (OH^–). Calculate the hydroxide ion concentration, [OH^–], in each well. Record the final concentrations of each ion in the Data Table under Ca²⁺ dilutions.
4. Repeat the calculations for the hydroxide ion and the calcium ion concentrations in the 12 wells with the hydroxide ion serial dilution. Record these final concentrations under OH^– dilutions.

Safety Precautions

Calcium nitrate solution is slightly toxic by ingestion. Sodium hydroxide solution is a corrosive liquid. Avoid contact with skin and especially the eyes. Wear chemical splash goggles, chemical-resistant gloves and a chemical-resistant apron. Wash hands thoroughly with soap and water before leaving the laboratory.

Procedure

Part I. Prepare a series of diluted calcium ion solutions.

1. Label one 50-mL beaker 0.10 M Ca(NO₃)₂ and another beaker 0.10 M NaOH. Place about 10 mL of the 0.10 M Ca(NO₃)₂ solution in the first beaker and 10 mL of 0.1 M NaOH in the other.
2. Place the reaction strip with 12 wells on a blank white sheet of paper for a good background.
3. Using a Beral microtip pipet, carefully add 5 drops of 0.10 M calcium nitrate solution in well 1. Note: Hold the Beral microtip pipet vertically when dispensing the drops. Make sure no air bubbles are in the pipet barrels. Discard the first drop into an empty waste beaker, as it may contain an air bubble. Empty the pipet into the empty waste beaker after adding the 5 drops.
4. Carefully add 5 drops of distilled water with a second Beral microtip pipet in each of the wells 2 through 12.
5. Using the empty pipet used in step 3, add 5 drops of 0.10 M calcium nitrate to well 2. Mix the solution thoroughly by stirring the solution with a plastic toothpick. Clean the toothpick with a clean paper towel. The solution in well 2 is now 0.050 M in Ca²⁺ ion. Make sure the pipet is empty before proceeding to the next step.
6. Using the Beral microtip pipet , remove the solution from well 2, and add 5 drops of this solution into well 3.
7. Put the remaining solution back in well 2. Note: Make sure all the liquid in the pipet transfers. Mix the solution in well 3 as in step 5.
8. Continue this serial dilution procedure, adding 5 drops of the previous solution to the 5 drops of water in each well down the row until the last one, well 12, is filled.
9. Mix the solution in well 12, and then discard 5 drops into the waste beaker. Each well should contain 5 drops of increasingly dilute Ca²⁺ ion solution.

Part II. Precipitation reaction.

1. With the third Beral pipet, carefully add 5 drops of 0.10 M sodium hydroxide solution, NaOH, into each of the wells 1 through 12. When the sodium hydroxide is added to each well, the initial concentrations of both reactants are halved, as each solution dilutes the other. The hydroxide ion concentration in each well is ½ x 0.10 M, or 0.05 M.
2. Use a clean toothpick to thoroughly mix each of these combined solutions. Wipe off the toothpick with a clean paper towel after each mixing.
3. Allow three or four minutes for the precipitates to form.
4. Observe the pattern of precipitation. In some of the wells the concentration of both ions is too low to have any precipitate form. Assume that the first well with no precipitate represents a saturated solution.
5. Record observations for each well in the Data Table in the Ca²⁺ dilutions row as either no precipitation, no ppt, or precipitation, ppt.
6. Rinse out the wells with three small portions of distilled or deionized water. Dry each well with a clean paper towel.

Part III. Check using serial dilutions of sodium hydroxide solution.

1. To check the results, repeat the previous procedure but use a serial dilution of the 0.10 M NaOH solution.
2. In the cleaned, dried 12-well reaction strip, put 5 drops of 0.10 M sodium hydroxide in well 1.
3. Put 5 drops of distilled water in wells 2 through 12.
4. Add 5 drops of the 0.10 M sodium hydroxide to well 2. Mix the solution by thoroughly stirring the solution with a clean toothpick. The solution in well 2 is now 0.050 M in OH^– ion.
5. Continue the serial dilution to well 12, and then remove 5 drops from well 12.
6. Add 5 drops of 0.10 M Ca(NO₃)₂ to each of the wells, and mix each with a clean toothpick. Again, determine the well where no more precipitate appears.
7. Record observations for each well in the Data Table in the OH^– dilutions row as either no precipitation, no ppt, or precipitation, ppt.
8. Dispose of the solutions as directed by the instructor.

Student Worksheet PDF

12980_Student1.pdf