Materials Included In Kit

Additional Materials Required

Prelab Preparation

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

• Iron(III) Nitrate, Fe(NO₃)₃, 0.0020 M: Using a volumetric pipet, transfer 5.00 mL of the standard 0.200 M ferric 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 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. Note: 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.

Safety Precautions

The iron(III) nitrate solution and the 1 M nitric acid are corrosive liquids; they 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 consult 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. The waste 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 one 50-minute class period. 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 solution in Part A. Using graduated cylinders to measure and transfer the liquids will not give the precision needed to achieve constant values of the equilibrium constant. Serological pipets 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 will save valuable time. Be sure to emphasize the need for students to thoroughly rinse the pipet before transferring each solution.
• The use of computer- or calculator-based technology for data collection and analysis is tailor-made for colorimetry experiments. The use of technology in this experiment also reflects the way technology is used in the “real world.” This is a key goal of technology integration in the curriculum. General instructions have been given in the Procedure section for electronic absorbance measurements. Absorbance measurements may also be made using a conventional spectrophotometer. A general procedure for the use of a conventional spectrophotometer is given in the Supplementary Information in the Further Extensions section.
• Students may well benefit from a review of the principles of light absorption and transmission (e.g., Why are some apples red?) before tackling colorimetry. All of the solutions in this experiment will 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 470 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 will absorb.
• 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.
• For more advanced classes, have students prepare a standard graph (usually called a calibration curve) of absorbance on the y-axis versus concentration on the x-axis for a series of reference solutions containing known amounts of FeSCN²⁺ ions. Some possible reference solutions are shown in the table below.
  {13950_Hints_Table_3}

  *Make up to 100.0 mL total volume in a 100-mL volumetric flask.

• Once the absorbance values of the reference solutions have been 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. A classic Beer’s Law plot of absorbance versus concentration has the form shown to the right. Alternatively, a calibration curve may be completed as part of a cooperative class project. Assign each group of students one of five different reference solutions to measure. Have students share results to average and graph the data. The points of absorbance versus concentration should fall on a straight line that goes through the origin. The unknown concentration of FeSCN²⁺ ions in each test solution can then be determined from the graph: find the absorbance value of the test solution, read across to the best-fit straight line through the data points and then down to the x-axis to find the concentration.
  {13950_Hints_Figure_2}
• Iron(III) nitrate is also called ferric nitrate. The iron(III) nitrate solutions must be prepared in nitric acid as the solvent because the reaction of Fe³⁺ and SCN^– ions is acid catalized.

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.

Further Extensions

Supplementary Information: 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 470 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 the reference solutions at 470 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 the reference solution in the data table.
3. Repeat step 2 for each of the test solutions. Record the absorbances in the 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.

Answers to Prelab Questions

1. The reference solution in Part A will be prepared by mixing 9.00 mL of 0.200 M Fe(NO₃)₃ solution and 1.00 mL of 0.0020 M KSCN solution. The concentration of Fe³⁺ ions in the reference solution (M₂) before any reaction occurs can be calculated using the so-called “dilution equation,” as shown.

   M₁V₁ = M₂V₂

   M₁ = concentration of solution before mixing = 0.200 M Fe(NO₃)₃
   V₁ = volume of solution before mixing = 9.00 mL
   V₂ = final volume of reference solution after mixing = 9.00 + 1.00 mL = 10.00 mL

   {13950_PreLab_Equation_5}
   Use the dilution equation to calculate the concentration of SCN– ions in the reference solution before any reaction occurs.

   M₁ = concentration of solution before mixing = 0.0020 M KSCN
   V₁ = volume of solution before mixing = 1.00 mL
   V₂ = volume of reference solution = 1.00 + 9.00 mL = 10.00 mL
   M₂ = M₁V₁/V₂ = (0.0020 M)(1.00 mL)/(10.00 mL) = 2.0 → 10^–4 M

2. “The equilibrium concentration of FeSCN²⁺ ions in the 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.

   The reference solution contains 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 thousandfold greater concentration of iron should be enough to consume essentially all of the thiocyanate ions and convert them to product.

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 data table. Hint: The final volume (V₂) of each test solution is 10.00 mL.
   {13950_PreLabAnswers_Table_4}

   Sample calculations for test solution 1:
   Concentration of Fe³⁺ ions = (0.0020 M)(5.00 mL)/(10.00 mL) = 1.0 → 10^–3 M
   Concentration of SCN^– ions = (0.0020 M)(1.00 mL)/(10.00 mL) = 2.0 → 10^–4 M
   The results are summarized on the next page.

   {13950_PreLabAnswers_Table_5}

Sample Data

Temperature ________________

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

Answers to Questions

1. As discussed in the Background section and the Prelab Questions, it is assumed that essentially all of the thiocyanate ions present in the reference solution will be converted to product. What is the concentration of FeSCN²⁺ ions in the reference solution?

   The concentration of FeSCN²⁺ ions in the reference solution is equal to 2.0 → 10^–4 M.

For Questions 2–7, construct a results table to summarize the results of the calculations.

2. For each test solution, the absorbance (A_n, where n = 1–5) should be directly proportional to the equilibrium concentration of FeSCN²⁺ ions. The concentration of FeSCN²⁺ ions can be calculated by comparing its absorbance versus that of the reference solution (A_ref). Use the following equation to calculate the equilibrium concentration of FeSCN²⁺ ions in each test solution 1–5. Enter the results in the results table.

   [FeSCN²⁺]_n = (A_n/A_ref) → [FeSCN²⁺]_ref

   Sample calculation for test solution 1:
   [FeSCN²⁺]₁ = (0.092/0.728) → (0.00020 M) = 2.5 → 10^–5 M

3. Calculate the equilibrium concentration of Fe³⁺ ions in each test solution 1–5: subtract the equilibrium concentration of FeSCN²⁺ ions from the initial concentration of Fe³⁺ ions (see the data table). Enter the results in the results table.

   [Fe³⁺]_eq,n = [Fe³⁺]_initial – [FeSCN²⁺]_n

   Sample calculation for test solution 1:
   [Fe³⁺]_eq,1 = 0.0010 M – 0.000025 M = 0.00098 M

4. Calculate the equilibrium concentration of SCN^– ions in each test solution 1–5: subtract the equilibrium concentration of FeSCN²⁺ ions from the initial concentration of SCN^– ions (see the data table). Enter the results in the results table.

   [SCN^–]_eq,n = [SCN^–]_initial – [FeSCN²⁺]_n

   Sample calculation for test solution 1:

   [SCN^–]_eq,1 = 0.00020 M – 0.000025 M = 0.00018 M

5. Use Equation 4 in the Background section to calculate the value of the equilibrium constant K_eq for each test solution 1–5. Enter the results in the results table.

   Sample calculation for test solution 1:

   {13950_Answers_Equation_7}
6. Calculate the mean (average value) of the equilibrium constant for the five test solutions.

   Mean = (140 + 180 + 150 + 140 + 140)/5 = 150

7. 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 1–5 is equal to the average deviation.

   Average deviation = (10 + 30 + 0 + 10 + 10)/5 = 12

   Results Table

   {13950_Answers_Table_7}
   Note: No units are given for the value of the equilibrium constant. This is the standard convention in most textbooks. Equilibrium constants are considered dimensionless quantities because they are defined thermodynamically using activities rather than concentrations.
8. 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 7% 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.

9. 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 cuvets, which are not precisely matched.

References

This experiment has been adapted from Flinn ChemTopic™ Labs, Volume 15, Equilibrium, Cesa, I., Ed., Flinn Scientific, Batavia, IL, 2003.

Introduction

For any reversible chemical reaction at equilibrium, the concentrations of all reactants and products are constant or stable. There will be 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 will eventually reach a position of chemical equilibrium. When this equilibrium occurs, the rate of the forward reaction just equals the rate of the reverse reaction. While both the forward and reverse reactions continue to occur, the net change in the concentrations of the reactants and products is zero and their equilibrium concentrations become constant.

Through an exhaustive number of experiments over the years, an empirical relationship has been found between the concentrations of the reactants and products of a reversible reaction at equilibrium. This relationship can be expressed mathematically in the form of the equilibrium constant, K_eq. 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. In some cases, equilibrium favors products and it appears that the reaction proceeds essentially to completion. The amount of reactants remaining under these conditions will be 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 will be very small.

The concentrations of reactants and products at equilibrium will vary, depending on the initial amounts of materials present. The special ratio of reactants and products described by K_eq will always be 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 to determine the equilibrium constant for 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 blood-red. The concentration of FeSCN²⁺ complex ions at equilibrium will be 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 will be. The wavelength of light absorbed by the red ions is about 470 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 will be directly proportional to the concentration of FeSCN²⁺ ions. When absorbance is plotted versus concentration of FeSCN²⁺ complex ions, a straight line relationship results (see Figure 1). Known concentrations of FeSCN²⁺ complex ions are used to create this curve. This curve is then used to determine the concentration of FeSCN²⁺ complex ions in any unknown sample.

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 will be 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 reference solution and a series of test solutions will be prepared. The reference solution will be prepared by mixing a large excess of Fe³⁺ ions with a known amount 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 solution will essentially be equal to the initial concentration of SCN^– ions. The test solutions will be prepared by mixing a constant concentration of Fe³⁺ ions with different concentrations of SCN^– ions. These solutions will contain unknown concentrations of FeSCN²⁺ ions at equilibrium.

In Part B, the absorbance of both the reference solution and the test solutions will be measured by colorimetry. The unknown concentrations of FeSCN²⁺ in the test solutions will be calculated by comparing their absorbance readings to the absorbance of the reference solution.

Materials

Prelab Questions

1. The reference solution in Part A will be prepared by mixing 9.00 mL of 0.200 M Fe(NO₃)₃ solution and 1.00 mL of 0.0020 M KSCN solution. The concentration of Fe³⁺ ions in the reference solution (M₂) before any reaction occurs can be calculated using the so-called “dilution equation,” as shown.

   M₁V₁ = M₂V₂

   M₁ = concentration of solution before mixing = 0.200 M Fe(NO₃)₃
   V₁ = volume of solution before mixing = 9.00 mL
   V₂ = final volume of reference solution after mixing = 9.00 + 1.00 mL = 10.00 mL

   {13950_PreLab_Equation_5}
   Use the dilution equation to calculate the concentration of SCN^– ions in the reference solution before any reaction occurs.
2. “The equilibrium concentration of FeSCN²⁺ ions in the 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.
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 data table. Hint: The final volume (V₂) of each test solution is 10.00 mL.
   {13950_PreLab_Table_1}

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 six 50-mL beakers or large test tubes and label them 1–6 for the test solutions and reference solution.
2. Rinsing the pipet before adding each reagent, combine the following volumes of reagents to prepare the test solutions. Note: There are two different iron “stock” solutions, 0.0020 M and 0.200 M Fe(NO₃)₃. Read the labels carefully before use!
   {13950_Procedure_Table_2}
3. Prepare the reference solution 6 by mixing 9.00 mL of 0.200 M Fe(NO₃)₃ and 1.00 mL of 0.0020 M KSCN in beaker 6. Note: Use the same pipet that was used in step 2.
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 data table. This will be assumed to be the equilibrium temperature for all of the solutions.

Part B. Colorimetry Measurements

6. Fill six cuvets about ¾-full with the solutions from Part A and arrange them in order on a labeled sheet of paper to keep track of the solutions. Do not write on the cuvets.
7. Handle the cuvets by their ribbed sides or their tops to avoid getting fingerprints on the surface. Wipe the cuvets with lint-free tissue paper or lens paper.
8. Connect the interface system to the computer or calculator and plug the colorimeter sensor into the interface.
9. Select Setup and Sensors from the main screen and choose “Colorimeter.” Note: Many newer sensors have an automatic calibration feature that automatically calibrates the colorimeter before use. If the sensor has the autocalibration feature, set the wavelength on the colorimeter to 470 nm (blue), press the autocalibration key, and proceed to step 14. If the sensor does not have the autocalibration feature, follow steps 10–13 to calibrate the colorimeter with a “blank” cuvet containing only distilled water.
10. Select Calibrate and Perform Now from the Experiment menu on the main screen.
11. Fill a cuvet about ¾-full with distilled water. Wipe the cuvet with a lint-free tissue, then place the cuvet in the colorimeter compartment.
12. Set the wavelength knob on the colorimeter to 0%T—the onscreen box should read zero. Press Keep when the voltage is steady.
13. Turn the wavelength knob on the colorimeter to 470 nm (blue)—the onscreen box should read 100. Press Keep when the voltage is steady.
14. Return to the main screen and set up a live readout and data table that will record absorbance (as a function of time).
15. Select Setup followed by Data Collection. Click on Selected Events to set the computer for manual sampling.
16. Remove the “blank” cuvet from the colorimeter compartment and replace it with the cuvet containing test solution 1.
17. Press Collect on the main screen to begin absorbance measurements.
18. When the absorbance reading stabilizes, press Keep on the main screen to automatically record the absorbance measurement. Note: The absorbance measurement should appear in a data table onscreen. The onscreen table will also contain a time reading, which may be ignored.
19. Remove the cuvet from the colorimeter compartment and replace it with the cuvet containing test solution 2.
20. When the absorbance reading stabilizes, press Keep on the main screen to automatically record the absorbance measurement.
21. Repeat steps 19 and 20 with the other test solutions 3–5 and with the reference solution 6.
22. Press Stop on the main screen to end the data collection process. If possible, obtain a printout of the data table.
23. Record the absorbance data for solutions 1–6 in the data table.
24. 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

13950_Student1.pdf