Lab #16: Determination of the Equilibrium Name: Constant of FeSCN 2+ Lab Exercise Chemistry II Partner: 10 points USE BLUE/BLACK INK!!!! Date: Hour: Goal: The goal of this lab is to determine the equilibrium constant of the reaction of iron (III) ions and thiocyanate ions forming the complex ion FeSCN 2+. Background: There are many reactions that take place in solution that are equilibrium reactions. They do not go to completion; both the forward and reverse reactions are occurring, and both reactants and products are always present. Examples of this type of reaction include weak acids, such as acetic acid, dissociating in water; weak bases, such as ammonia, reacting with water; and the formation of complex ions in which a metal ion combines with one or more negative ions. In this lab, a reaction involving the formation of complex ions from solutions of iron (III) and thiocyanate ions will be studied, and the equilibrium constant of the reaction will be determined. Chemical reactions are driven to completion by two forces: a decrease in energy (exothermic reactions), or an increase in entropy. If both an energy decrease and an entropy increase occur in the forward reaction, the reaction will go to completion. An example of this type of reaction is combustion the reaction is exothermic and has an increase in entropy, so it goes to completion. However, when an energy decrease drives a reaction in one direction and an entropy increase drives it in the reverse direction, equilibrium will result. The reaction will not go to completion, but it will reach a point where both reactants and products are present in a fixed ratio of concentrations. The reaction will continue at the same rate in both forward and reverse directions, and the concentrations of products and reactants will stay constant. These ideas can be expressed mathematically in the form of the equilibrium constant. Consider the following general equation for a reversible chemical reaction: aa + bb W cc + dd The equilibrium constant, K eq, for this general equation is K eq c C D = [ ] [ ] a [ A] [ B] d b, where 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 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 products to reactants described by the K eq is always the same 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 if iron (III) ions with thiocyanate ions provides a convenient example for determining the equilibrium constant of a reaction. Fe 3+ and SCN - ions combine to form the FeSCN 2+ ion according to the equation Fe 3+ (aq) + SCN - (aq) W FeSCN 2+ (aq) pale yellow colorless blood-red The value of K eq can be determined experimentally by mixing known concentrations of Fe 3+ and SCN - ions and measuring the concentration of FeSCN 2+ ions at equilibrium. The concentration of FeSCN 2+ complex ions at equilibrium is proportional to the intensity of the red color. A spectrophotometer will be used to measure the concentration of FeSCN 2+ ions using Beer s Law (also called the Beer-Lambert Law), A = g b c where A is the absorbance, g is the absorptivity of the sample, b is the cell path length, and c is the molar concentration of the substance absorbing the light. A graph of absorption verses concentration is a straight line. This reaction has two parts. In the first part, a series of reference solutions and test solutions are prepared. The reference solutions are prepared by mixing a large excess of Fe 3+ ions with known, smaller amounts of SCN - ions. According to Le Châtelier s Principle, the large excess of iron (III) ions should effectively convert all of the thiocyanate ions to the blood-red FeSCN 2+ complex ions. The concentration of FeSCN 2+ 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 Fe 3+ ions with different amounts of SCN - ions. These solutions contain unknown concentrations of FeSCN 2+ ions at equilibrium. In the second part, the absorbances of both the reference solutions and the test solutions are measured by a spectrophotometer. A calibration curve is constructed from the absorption values of the reference solutions. The unknown concentrations of FeSCN 2+ in the test solutions are calculated by comparing their absorbances to the absorbance values on the calibration curve. These values are then used to determine the equilibrium concentrations and the equilibrium constant for the reaction. Research questions (Please answer on a separate sheet and attach): 1) Define chemical equilibrium. 2) Briefly explain Le Châtelier s Principle. 3) A similar reaction to our is Ag + (aq) + 2 NH 3 (aq) W Ag(NH 3 ) 2 + (aq) (a) Write the equilibrium constant expression for the reaction. (b) An experiment was carried out to determine the value of the K eq for the reaction. The following data were collected: Initial moles of Ag + = Initial moles of NH 3 present = 3.6 x 10-3 mol 6.9 x 10-3 mol
+ Measured concentration of Ag(NH 3 ) 2 at equilibrium = 3.4 x 10-2 M Total volume of solution 100. ml Calculate the number of unreacted moles of Ag + at equilibrium. Calculate the molarity of the unreacted Ag + at equilibrium. Calculate the number of unreacted moles of NH 3 at equilibrium. Calculate the molarity of the unreacted NH 3 at equilibrium. Use the molarities to calculate the value of the equilibrium constant. 4) Use the dilution equation V 1 M 1 = V 2 M 2 to calculate the concentration of the SCN - ions in the five reference solutions before any reaction occurs. Enter these values in Data Table 1 as [FeSCN 2+ ]. 5) Use the same dilution equation to calculate the concentrations of Fe 3+ and SCN - ions in each test solution after mixing them together but before any reaction occurs. Enter the results of these calculations in scientific notation in Data Table 2. Materials: 2 burets 50 ml 0.200 M Fe(NO 3 ) 3 in 1 M HNO 3 1 double buret clamp 35 ml 0.0020 M Fe(NO 3 ) 3 in 1 M HNO 3 1 ringstand 25 ml 0.0020 M KSCN 30 ml 0.00020 M KSCN 10 mm test tubes 1 test tube brush 4 50 ml beakers 1 permanent marker 6 cuvets 1 spectrophotometer 1 thermometer 2 buret funnels 1 glass stirring rod 10 # solid rubber stoppers 2 test tube racks 1 Kimwipe Hazards:
Procedure - Day 1: 1) Physically and chemically clean the beakers, stirring rod, test tubes, buret funnels, and burets. 2) Obtain the four solutions in separate, labeled beakers. 3) Prepare the reference solutions a) Fill one buret with 40 ml of 0.200 M Fe(NO 3 ) 3 solution. b) Label the 5 of the test tubes Reference Soln 1-5 and initials c) Add the listed amount of 0.200 M Fe(NO 3 ) 3 solution to each of the test tubes Sample Volume of 0.200 M Fe(NO 3 ) 3 solution Volume of 0.00020 M KSCN solution Reference sol n #1 8.0 ml 2.0 ml Reference sol n #2 7.0 ml 3.0 ml Reference sol n #3 6.0 ml 4.0 ml Reference sol n #4 5.0 ml 5.0 ml Reference sol n #5 4.0 ml 6.0 ml d) Fill the second buret with 30 ml of 0.00020 M KSCN solution e) Add the listed amount of 0.00020 M KSCN solution to each of the test tubes f) Stir each test tube with the stirring rod, cleaning and drying it between each to prevent cross-contamination. 4) Re-clean the burets 5) Prepare the test solutions a) Fill the first buret with 35 ml of 0.0020 M Fe(NO 3 ) 3 solution b) Label 5 more of the test tubes Test Samples 6-10 and initials c) Dispense 5.00 ml of 0.0020 M Fe(NO 3 ) 3 solution into each test tube d) Fill the second buret with 25 ml of 0.0020 M KSCN solution e) Add the listed amount of 0.0020 M KSCN solution to each of the test tubes Sample Volume of 0.0020 M Fe(NO 3 ) 3 solution Volume of 0.0020 M KSCN solution Volume of Distilled Water Added Test sol n #6 5.0 ml 1.0 ml 4.0 ml Test sol n #7 5.0 ml 2.0 ml 3.0 ml Test sol n #8 5.0 ml 3.0 ml 2.0 ml Test sol n #9 5.0 ml 4.0 ml 1.0 ml Test sol n #10 5.0 ml 5.0 ml 0.0 ml f) Stir each test tube with the stirring rod, cleaning and drying it between each to prevent cross-contamination.
6) Stopper each labeled test tube 7) Clean up Procedure - Day 2: 1) Measure the temperature of one of the solutions. (This will be the equilibrium temperature.) 2) Turn on and warm up the spectrophotometer for at least 15 minutes 3) Physically and chemically clean the cuvets 3) Set the spectrophotometer at 450 nm 4) With nothing in the spectrophotometer, set the transmittance to zero. 5) Fill one cuvet with distilled water. (This will be the blank) 6) With the distilled water cuvet in the spectrophotometer, set the transmittance to 100%. 7) Measure the transmittance of each of the reference solutions at 450 nm. 8) Measure the transmittance of each of the test solutions at 450 nm. 9) Clean up! Data: Data Table 1: Reference Solutions Sample [FeSCN 2+ ] % T Reference sol n #1 Reference sol n #2 Reference sol n #3 Reference sol n #4 Reference sol n #5 Data Table 2: Test Solutions Sample [Fe 3+ ] just after mixing [SCN - ] just after mixing %T Test sol n #6 Test sol n #7 Test sol n #8 Test sol n #9 Test sol n #10
Post-Lab Calculations, Analysis, and Questions (answer on separate paper and/or the computer no need for a separate conclusion): 1. Convert % T to absorbance using the formula 2 - %T. 2. Plot molar concentration of FeSCN 2+ versus absorbance. Put a best-fit straight line ( trendline in Excel terms) through the data points. Include the origin as a valid point. 3. Determine the unknown concentration of FeSCN 2+ in each test solution from the graph. 4. Make a results table in Excel (or on paper) with this information: Sample [FeSCN 2+ ] eq [Fe 3+ ] eq [SCN - ] eq K eq Test sol n #6 Test sol n #7 Test sol n #8 Test sol n #9 Test sol n #10 Average value Average deviation 5. Record the FeSCN 2+ concentration for each test solution in the results table. 6. Calculate the equilibrium concentration of Fe 3+ ions in each test solution by subtracting the equilibrium concentration of FeSCN 2+ ions from the initial concentration of Fe 3+ ions (from the test solutions data table). [Fe 3+ ] eq = [Fe 3+ ] at mixing - [FeSCN 2+ ] eq 7. Calculate the equilibrium concentration of SCN - ions in each test solution by subtracting the equilibrium concentration of FeSCN 2+ ions from the initial concentration of SCN - ions (from the test solutions data table). [SCN - ] eq = [SCN-] at mixing - [FeSCN 2+ ] eq 8. Calculate the equilibrium constant for each test solution. 9. Calculate the mean (average) value of the equilibrium constant for the five test solutions. 10. Calculate the average deviation for the equilibrium constant by finding the absolute value of the difference between each individual value of the equilibrium constant and the average. The average of the differences is the average deviation. (This is your precision or uncertainty.) 11. Was your equilibrium constant actually constant? Should it have been constant? Explain your answer. 12. What does the calculated equilibrium constant indicate about the degree of completion of the reaction? At equilibrium, are there mostly products, reactants, or both? 13. Explain how the spectrophotometer worked during this lab and how it was used to determine the answer. 14. What measurement(s) limited the number of significant figures in this lab? 15. What were the problems in this lab? How could this lab be improved? Lab handout based on the experiment The Determination of K eq for FeSCN 2+ in Laboratory Experiments for Advanced Placement Chemistry (Second Edition) by S.A. Vonderbrink (Flinn Scientific, 2006)