Materials Included In Kit

Additional Materials Required

Prelab Preparation

For best results, prepare all of the solutions with analytical precision using an analytical balance and volumetric flasks, as described.

Iron(III) Nitrate, Fe(NO₃)₃, 0.200 M: Add 40.40 g of iron(III) nitrate nonahydrate [Fe(NO₃)₃•9H₂O] to about 100 mL of 1 M nitric acid in a 500-mL volumetric flask. Mix thoroughly to dissolve, then dilute to the mark with 1 M nitric acid. Mix well prior to dispensing.

Iron(III) Nitrate, Fe(NO₃)₃, 0.0020 M: Using a volumetric pipet, transfer 5.00 mL of the standard 0.200 M iron(III) nitrate solution to a 500-mL volumetric flask half-filled with 1 M nitric acid. Dilute to the mark with 1 M nitric acid and mix well prior to dispensing.

Potassium Thiocyanate, KSCN, 0.0020 M: Dissolve 0.097 g of potassium thiocyanate (KSCN) in about 250 mL of distilled or deionized water in a 500-mL volumetric flask. Mix thoroughly to dissolve, then dilute to the mark with distilled water. Mix well prior to dispensing. Do not use 1 M nitric acid as the solvent for this solution. Thiocyanate ions decompose in the presence of nitric acid. Note: If an analytical balance is not available, prepare a more concentrated solution (for example, a 0.200 or 0.020 M solution). Then dilute the more concentrated solution using a pipet and volumetric flask.

Potassium Thiocyanate, KSCN, 0.00020 M: Using a buret or volumetric pipet, transfer 50.0 mL of the 0.0020 M KSCN solution to a 500-mL volumetric flask. Dilute to the mark with distilled water. Prepare this solution daily before the experiment.

Safety Precautions

Iron(III) nitrate solution contains 1 M nitric acid and is a corrosive liquid; it will stain skin and clothing. Instruct students to notify the teacher immediately in case of a spill. Keep sodium carbonate or sodium bicarbonate on hand to clean up spills. Potassium thiocyanate is toxic by ingestion; it can generate poisonous hydrogen cyanide gas if heated strongly. Avoid contact of all chemicals with eyes and skin. Wear chemical splash goggles and chemical-resistant gloves and apron. Wash hands thoroughly with soap and water before leaving the laboratory. 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 regulation that may apply, before proceeding. The leftover solutions may be combined and neutralized and then flushed down the drain with excess water according to Flinn Suggested Disposal Method #24b.

Lab Hints

• The laboratory work for this experiment can reasonably be completed in two 50-minute class periods. The Prelab Questions may be assigned separately as preparation for lab, or they may be used as the basis of a cooperative class discussion.
• Mohr or serological-type pipets are recommended for preparation of the test solutions and the reference solutions in Part A. Using graduated cylinders to measure and transfer the liquids does not give the precision needed to achieve constant values of the equilibrium constant. Serological pipets (see Flinn Catalog No. GP7059) are considered “throw-away” pipets by the medical community and are very affordable. They can be reused several times before the graduations come off. Used as disposable pipets, however, they eliminate the need for dishwashing and save valuable time.
• The use of computer- or calculator-based technology for data collection and analysis is tailormade for colorimetry experiments.
• Nitric acid is used instead of either hydrochloric acid or sulfuric acid to acidify the iron(III) solution. Hydrochloric acid and sulfuric acid impart a yellow color to the iron(III) solution that interferes with the absorbance values of the reference solutions.
• The precise concentration of nitric acid in the 0.200 M or 0.0020 M solutions of iron(III) nitrate is not critical. The nominal amount of 1 M nitric included with the kit is 1000 mL, enough to fill two 500-mL volumetric flasks exactly to the mark to make the required analytical Fe(NO₃)₃ solutions, as described in the Preparation of Solutions section. However, the actual volume of nitric acid may deviate from this stated amount by 2-6 mL. Please fill each 500-mL volumetric flask with 1 M nitric acid included with the kit, and then add 1–2 mL distilled or deionized water, if needed, to fill each flask to the mark. As a result of this action, the concentration of nitric acid may vary from a nominal value of 1.0 M to 0.996 M. This difference is not significant and will not affect the performance of the lab or the accuracy of the results, as long as the iron(III) concentration is precisely known. The acid is needed to make sure the solutions are acidic. The precise concentration of acid is not critical.
• Students may well benefit from a review of the principles of light absorption and transmission— why are some apples red—before tackling colorimetry. All of the solutions in this experiment have a red color. That means that when ordinary “white” light is passed through the solution, only the red color (wavelength) is transmitted. All of the other colors (wavelengths of light) are absorbed. The colorimeter sends a beam of monochromatic (one color or wavelength) light through the solution. The wavelength of light used in this experiment is 450 nm, corresponding to blue light. This is the complementary color of the red color of the solution and is thus almost totally absorbed (not transmitted) by the solution. The beam of light is passed through the sample, and the intensity of the blue light that is transmitted is measured electronically. The rest of the blue light is absorbed by the solution. The greater the concentration of red FeSCN²⁺ ions in solution, the more blue light the solution absorbs.
• The “Color and Light Spectrum Demonstrations Kit” available from Flinn Scientific (Catalog No. AP6172) uses a holographic diffraction grating and an overhead projector to produce a very large visible spectrum. Placing a beaker of red water in the “slit” opening clearly shows that all of the colors except red are absorbed by the red solution.
• Once the absorbance values of the reference solutions are measured, a calibration curve may be obtained automatically using the computer interface system and software to generate a Beer’s Law plot of absorbance versus concentration.

Teacher Tips

• The term “equilibrium position” may be confusing to students. It is difficult to reconcile the ideas that there are an infinite number of equilibrium positions but only one unique equilibrium constant for a given reaction at a fixed temperature. It may be helpful to lead into the experiment with a visual introduction. Prepare samples of the equilibrium test solutions. Show students the solutions and ask them to discuss what is the same and what is different in these solutions. All of the solutions contain both reactants and products. The reactants and products are present in different concentrations in each solution. All solutions, however, are at equilibrium. The equilibrium constant is a characteristic property that should be the same for all of the solutions.
• Another difficulty with equilibrium is that students think that once equilibrium is reached, the reaction stops. This is incorrect. The forward and reverse reactions continue but the total amounts of reactants and products do not change.

Further Extensions

AP^® Standards
This laboratory fulfills the requirements for the College Board recommended AP Experiment #17: Colorimetric or Spectrophotometric Analysis. In addition, this lab also fulfills the requirements for AP Experiment #10: Determination of the Equilibrium Constant for a Chemical Reaction.

Answers to Prelab Questions

1. “The equilibrium concentration of FeSCN²⁺ ions in the reference solutions is essentially equal to the concentration of SCN^– ions in solution before any reaction occurs.” Use Le Chatelier’s Principle to explain why this statement is true.

   The reference solutions contain a large excess of iron(III) ions relative to the concentration of thiocyanate ions. According to Le Chatelier’s Principle, adding more reactants should shift the equilibrium to the right, that is, to make more product. The “stress” on the reactant side due to the presence of a thousand-fold greater concentration of iron should be enough to consume essentially all of the thiocyanate ions and convert them to product.

2. The five reference solutions in Part A are prepared by mixing the 0.200 M Fe(NO₃)₃ solution and the 0.00020 M KSCN solution in the amounts listed in the following table.
   {13858_PreLabAnswers_Table_5}
   The concentration of Fe³⁺ ions in the first reference solution (M₂) before any reaction occurs can be calculated using the so-called “dilution equation,” as shown
   {13858_PreLabAnswers_Equation_5}

   M₁ = concentration of solution before mixing = 0.200 M Fe(NO₃)₃
   V₁ = volume of solution before mixing = 8.0 mL
   V₂ = final volume of reference solution after mixing = 8.0 + 2.0 mL = 10.0 mL

   {13858_PreLabAnswers_Equation_6}
   Use the dilution equation to calculate the concentration of SCN^– ions in the five reference solutions before any reaction occurs. Enter these values in the Reference Solutions Data Table as [FeSCN²⁺].

   Reference Solution 1

   M₁ = concentration of solution before mixing = 0.00020 M KSCN
   V₁ = volume of solution before mixing = 2.0 mL
   V₂ = final volume of reference solution after mixing = 2.0 + 8.0 mL = 10.0 mL

   {13858_PreLabAnswers_Equation_7}

   Reference Solution 2

   M₂ = 0.000060 M or 6.0 x 10^–5 M

   Reference Solution 3

   M₂ = 0.000080 M or 8.0 x 10^–5 M

   Reference Solution 4

   M₂ = 0.00010 M or 1.0 x 10^–4 M

   Reference Solution 5

   M₂ = 0.00012 M or 1.2 x 10^–4 M

3. The following table summarizes the volumes of Fe³⁺ and SCN^– stock solutions that are mixed together to prepare the test solutions in Part A. Use the dilution equation to calculate the concentrations of Fe³⁺ and SCN^– ions in each test solution before any reaction occurs. Enter the results of these calculations in scientific notation in the Test Solutions Data Table. Hint: The final volume (V₂) of each test solution is 10.0 mL.
   {13858_PreLabAnswers_Table_6}
   Sample calculations for test solution 6: Concentration of Fe³⁺ ions = (0.0020 M)(5.0 mL)/(10.0 mL) = 1.0 x 10^–3 M Concentration of SCN^– ions = (0.0020 M)(1.0 mL)/(10.0 mL) = 2.0 x 10^–4 M These calculations are summarized in the table.
   {13858_PreLabAnswers_Table_7}

Sample Data

Reference Solutions

Test Solutions

*These are the concentrations of ions in solution immediately after mixing and before any reaction has occurred. See the Prelab Questions for calculations.

Results Table

Answers to Questions

1. Calculate the equilibrium concentration of Fe³⁺ ions in each test solution 6–10: Subtract the equilibrium concentration of FeSCN²⁺ ions from the initial concentration of Fe³⁺ ions (see the Test Solutions Data Table). Enter the results in the Results Table.

   [Fe³⁺]eq = [Fe³⁺]initial – [FeSCN²⁺]_eq

   Sample calculation for test solution 6:
   [Fe³⁺]_eq = 0.0010 M – 0.000027 M = 0.00097 M

2. Calculate the equilibrium concentration of SCN^– ions in each test solution 6–10: subtract the equilibrium concentration of FeSCN²⁺ ions from the initial concentration of SCN^– ions (see the Test Solution Data Table). Enter the results in the Results Table.

   [SCN^–]_eq = [SCN^–]_initial – [FeSCN²⁺]_eq

   Sample calculation for test solution 6:
   [SCN^–]_eq = 0.00020 M – 0.000027 M = 0.00017 M

3. Use Equation 4 in the Background section to calculate the value of the equilibrium constant K_eq for each test solution 6–10. Enter the results in the Results Table.

   Sample calculation for test solution 6:

   {13858_Answers_Equation_8}
4. Calculate the mean (average value) of the equilibrium constant for the five test solutions.

   Mean = (160 + 140 + 150 + 150 + 160)/5 = 150

5. Calculate the average deviation for K_eq: Find the absolute value of the difference between each individual value of the equilibrium constant and the mean. The average of these differences for solutions 6–10 is equal to the average deviation.

   Average deviation = (10 + 10 + 0 + 0 + 10)/5 = 8

6. The average deviation describes the precision of the results. Does the precision indicate that the equilibrium constant is indeed a “constant” for this reaction? Explain.

   The average deviation represents a 5.3% uncertainty in the equilibrium constant. Since there were only two significant figures in the concentration and volume measurements, the precision of the results would seem to fall within the range of uncertainty in the measurements themselves. Within the limits of experimental error, therefore, it appears that the equilibrium constant is indeed a constant.

7. Describe the possible sources of error in this experiment and their likely effect on the results.

   Note: Many students are likely to focus their discussion here on technique errors involving measuring and transferring liquids. In this experiment, these errors are probably very significant. The procedure requires considerable skill in using pipets. Technique errors will lead to poor precision in the results—the equilibrium constant will not look like a constant. Another source of error is that absorbance measurements are made in different cuvettes, which are not precisely matched.

Introduction

For any reversible chemical reaction at equilibrium, the concentrations of all reactants and products are constant or stable. There is no further net change in the amounts of reactants and products unless the reaction mixture is disturbed in some way. The equilibrium constant provides a mathematical description of the position of equilibrium for any reversible chemical reaction. What is the equilibrium constant and how can it be determined?

Concepts

• Chemical equilibrium
• Equilibrium constant
• Complex-ion reaction
• Colorimetry

Background

Any reversible reaction eventually reaches a position of chemical equilibrium. In some cases, equilibrium favors products and it appears that the reaction proceeds essentially to completion. The amount of reactants remaining under these conditions is very small. In other cases, equilibrium favors reactants and it appears that the reaction occurs only to a slight extent. Under these conditions, the amount of products present at equilibrium is very small.

These ideas can be expressed mathematically in the form of the equilibrium constant. Consider the following general equation for a reversible chemical reaction:

The equilibrium constant K_eq for this general reaction is given by Equation 2, where the square brackets refer to the molar concentrations of the reactants and products at equilibrium. The equilibrium constant gets its name from the fact that for any reversible chemical reaction, the value of K_eq is a constant at a particular temperature. The concentrations of reactants and products at equilibrium vary, depending on the initial amounts of materials present. The special ratio of reactants and products described by K_eq is always the same, however, as long as the system has reached equilibrium and the temperature does not change. The value of K_eq can be calculated if the concentrations of reactants and products at equilibrium are known.

The reversible chemical reaction of iron(III) ions (Fe³⁺) with thiocyanate ions (SCN^–) provides a convenient example for determining the equilibrium constant of a reaction. As shown in Equation 3, Fe³⁺ and SCN^– ions combine to form a special type of combined or “complex” ion having the formula FeSCN²⁺. The equilibrium constant expression for this reaction is given in Equation 4. The value of K_eq can be determined experimentally by mixing known concentrations of Fe³⁺ and SCN^– ions and measuring the concentration of FeSCN²⁺ ions at equilibrium. As noted in Equation 3, the reactant ions are pale yellow and colorless, respectively, while the product ions are bloodred. The concentration of FeSCN²⁺ complex ions at equilibrium is proportional to the intensity of the red color.

A special sensor or instrument called a colorimeter can be used to measure the absorbance of light by the red ions. The more intense the red color, the greater the absorbance. The wavelength of light absorbed by the red ions is about 450 nm. None of the other ions present in solution absorb light at this wavelength. As long as the same size container is used to measure the absorbance of each solution, the absorbance is directly proportional to the concentration of FeSCN²⁺ ions.

Experiment Overview

The purpose of this experiment is to calculate the equilibrium constant for the reaction of iron(III) ions with thiocyanate ions. The reaction is tested under different conditions to determine if the equilibrium constant always has the same numerical value. There are two parts to the experiment.

In Part A, a series of reference solutions and test solutions are prepared. The reference solutions are prepared by mixing a large excess of Fe³⁺ ions with known amounts of SCN^– ions. According to Le Chatelier’s Principle, the large excess of iron(III) ions should effectively convert all of the thiocyanate ions to the blood-red FeSCN²⁺ complex ions. The concentration of FeSCN²⁺ complex ions in the reference solutions is essentially equal to the initial concentration of SCN^– ions. The test solutions are prepared by mixing a constant amount of Fe³⁺ ions with different amounts of SCN^– ions. These solutions contain unknown concentrations of FeSCN²⁺ ions at equilibrium.

In Part B, the absorbances of both the reference solutions and the test solutions are measured by colorimetry. A calibration curve is constructed from the absorption values of the reference solutions. The unknown concentrations of FeSCN²⁺ in the test solutions are calculated by comparing their absorbance readings to the absorbance values of the calibration curve.

Materials

Prelab Questions

1. “The equilibrium concentration of FeSCN²⁺ ions in each reference solution is essentially equal to the concentration of SCN^– ions in solution before any reaction occurs.” Use Le Chatelier’s Principle to explain why this statement is true.
2. The five reference solutions in Part A are prepared by mixing the 0.200 M Fe(NO₃)₃ solution and the 0.00020 M KSCN solution in the amounts listed in the following table.
   {13858_PreLab_Table_1}
   The concentration of Fe³⁺ ions in the first reference solution (M₂) before any reaction occurs can be calculated using the so-called “dilution equation,” as shown.
   {13858_PreLab_Equation_5}
   M₁ = concentration of solution before mixing = 0.200 M Fe(NO₃)₃ V₁ = volume of solution before mixing = 8.0 mL V₂ = final volume of reference solution after mixing = 8.0 + 2.0 mL = 10.0 mL
   {13858_PreLab_Equation_6}
   Use the dilution equation to calculate the concentration of SCN^– ions in the five reference solutions before any reaction occurs. Enter these values in the Reference Solutions Data Table as [FeSCN²⁺].
3. The following table summarizes the volumes of Fe³⁺ and SCN^– stock solutions that will be mixed together to prepare the test solutions in Part A. Use the dilution equation to calculate the concentrations of Fe³⁺ and SCN^– ions in each test solution before any reaction occurs. Enter the results of these calculations in scientific notation in the Test Solutions Data Table. Hint: The final volume (V₂) of each test solution is 10.0 mL.
   {13858_PreLab_Table_2}

Safety Precautions

Iron(III) nitrate solution contains 1 M nitric acid and is a corrosive liquid; it will stain skin and clothing. Notify the teacher and clean up all spills immediately. Potassium thiocyanate is toxic by ingestion; it can generate poisonous hydrogen cyanide gas if heated strongly. Avoid contact of all chemicals with eyes and skin. Wear chemical splash goggles and chemical-resistant gloves and apron. Wash hands thoroughly with soap and water before leaving the laboratory.

Procedure

Part A. Preparing the Solutions

1. Obtain ten 50-mL beakers or large test tubes.
2. Prepare the five reference solution test tubes or beakers listed in the table below. Use a separate pipet to transfer the appropriate volumes of each reagent. Mix each solution using a stirring rod. Rinse the stirring rod and dry it between solutions. Label the test tubes or beakers with the corresponding reference solution number.
   {13858_Procedure_Table_3}
3. Using a separate pipet for each reagent to be added, combine the following volumes of reagents to prepare the test solutions. Note: Label the tubes with the corresponsing solution numbers 6 through 10. Read the reagent labels carefully before use!
   {13858_Procedure_Table_4}
4. Mix each solution using a stirring rod. Rinse the stirring rod and dry it between solutions.
5. Measure the temperature of one of the solutions and record it in the Test Solutions Data Table. This is assumed to be the equilibrium temperature for all of the solutions.

Part B. Colorimetry Measurements

1. Follow the procedure for your colorimetric measurements of the solution as directed by the instructor. Generally, spectrophotometers are used as follows: Turn the instrument on and allow it to warm up for 15 minutes. Set the wavelength at 450 nm. With no light passing through the instrument to the phototube, set the percent transmittance to zero with the “zero” control. Handle cuvets at the top so no fingerprints are in the light path. Polish cuvets with a tissue. Place a cuvet which is about ⅔ full of distilled water into the sample holder and set the percent transmittance to 100% with the appropriate control (not the zero control). Fill a cuvet about ⅔ full of a test solution, place it in the spectrophotometer and read the absorbance. Consult the instrument manual for details on its use.
2. Measure the absorbance of each of the reference solutions at 450 nm, using distilled water as the zero absorbance reference in the spectrophotometer. If absorbance is difficult to measure precisely on the meter because it is in the high range where the numbers are close together, measure percent transmittance and calculate the absorbance for each solution. Absorbance = –log T, where T is transmittance expressed as a decimal. Record the absorbance value for each reference solution used in the Reference Solutions Data Table.
3. Repeat step 2 for each of the test solutions. Record the absorbances in the Test Solution Data Table.
4. Dispose of the contents of the cuvets and of the remaining test solutions as directed by your instructor. Follow your instructor’s directions for rinsing and drying the cuvets.

Student Worksheet PDF

13858_Student1.pdf