K = [C]c [D] d [A] a [B] b (5)

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1 Chem 1B Dr. White 19 Experiment 3: Determination of an Equilibrium Constant Objectives To determine the equilibrium constant for a reaction. Introduction Equilibrium is a dynamic state in which, at a given temperature, a chemical reaction will reach a point where the molar ratio of products to reactants (otherwise known as the equilibrium constant) reaches a constant value. Another way to describe this is that the rate of formation of products is equal to the rate of formation of the reactants. Reactions that have reached equilibrium will remain so unless the system is disturbed by some outside factor. The reaction will then minimize the stress and restore the ratio of reactants to products. For the general reaction: aa + bb cc + dd (4) the equilbirum constant is described by: K = [C]c [D] d [A] a [B] b (5) where the square brackets ([ ]) indicate concentration in units of Molarity (M or mole/l). However, the equilbrium constant K is always unitless. At a given temperature, a chemical reaction will always reach a state of equilibrium according to its given equilibrium constant, K, no matter how the reaction is performed (i.e. A, B, C, and D are mixed in arbitrary amounts). To demonstrate this phenomenon, we will study the chemical reaction of iron(iii) ion and thiocyanate ion (SCN - ) shown below. Fe 3+ (aq) + SCN - (aq) FeSCN 2+ (aq) (6) When solutions containing the colorless iron (III) ion and colorless thiocyanate ion are mixed, the thiocyanatioiron (III) ion product is formed which is a deep red color. By determining the concentrations of these three chemical species in several solutions, we can calculate the equilibrium constant K for this reaction. At constant temperature the value of K is a constant, however the value of K will actually change somewhat with reactant concentration because of a complicated side reaction where the deep red thiocyantoiron(iii) ion product will also react with an additional thiocyanate ion to form a dithiocyanatoiron(iii) ion. Today, standard solutions of iron(iii) nitrate and thiocyanic acid will be used to prepare the equilibrium sample solutions. The amount of hydrogen ion will be kept approximately constant in all solutions using nitric acid. The concentration of the colored thiocyanatoiron (III) ion product will be measured by absorption spectrometry. Using this information and the Beer- Lambert plot from experiment 2, the concentrations of each species at equilibrium can be determined and the value of K can also be calculated. Tube # Procedure 1. Prepare a blank solution (Solution 0) and three equilibrium solutions according to the table below into four clean, dry test tubes using pipets to measure the volumes: Volume of M Fe(NO 3) 3 (ml) Volume of M KSCN (ml) Volume of HNO 3 (ml) Vortex each tube to mix thoroughly. 3. Calibrate the spectrometer by filling a clean, dry cuvette ¾ full with Solution 0.. Use a grease pencil to make a small mark on the cuvette so that you can place it in the same orientation throughout the experiment. Wipe the cuvette with a KimWipe and place it in the spectrometer. Select Calibrate Spectrometer from the Experiment menu. The calibration dialog box will display the message: Waiting... seconds for lamp to warm up. The minimum warmup time is one minute. Follow the instructions in the dialog box to complete the calibration. Click. 4. Empty the cuvette and rinse it twice with small amounts of the solution 1. Fill the cuvette ¾ full with solution 1 and place it in the spectrometer. Click the Configure Spectrometer Data Collection

2 20 Chem 1B Dr. White icon,, on the toolbar. A dialog box will appear. Select Abs vs. Wavelength under Set Collection Mode. Click to proceed. 5. Click. Click to complete the analysis. Record the absorbance value at the same λ max from Experiment 2. Complete the procedure for solutions 2-3. Be sure to record the absorbance values in your lab notebook. 6. Using your absorbance values measured for each tube in today s lab and your Beer- Lambert Plot from Experiment 2, determine the equilibrium concentration of FeSCN 2+ (the value of "x" in your "ICE" table below). Since FeSCN 2+ was generated by the reaction of Fe 3+ and SCN 1-, the amount of Fe 3+ and SCN 1- each diminished by the amount of FeSCN 2+ formed. To determine the equilibrium initial concentrations of Fe 3+, and SCN 1- for each tube, complete the following "ICE" table for each tube in your lab notebook. Fe 3+ SCN 1- FeSCN 2+ Calculate the conc. here based on 2.50 ml of M Fe 3+ in 5.00 ml total Calculate the conc. here based on x ml of M KSCN in 5.00 ml total change - x - x + x equilibrium [Fe3+ ] initial - x [SCN 1- ] initial - x 0 since no FeSCN 2+ was initially added to the tubes 7. Using the equilibrium concentrations of Fe 3+, SCN -, and FeSCN 2+, determine the value of K for each tube and determine an average K. x

3 Chem 1B Dr. White 21 Name: Lab Day/Time: The Fe 3+ + SCN - FeSCN 2+ Equilibrium System Initial Concentrations of FeSCN 2+, Fe 3+ and SCN - : Experiment 3 Determination of an Equilibrium Constant Data and Results Solution# [Fe 3+ ] initial (mol/l) [SCN - ] initial (mol/l) [FeSCN +2 ] initial (mol/l) Show calculations for how you obtained the Fe 3+ and SCN - initial concentrations in Solution 1 (as calculated in the OWL Pre-lab assignment). Equilibrium Concentrations of FeSCN 2+, Fe 3+ and SCN - : Solution # 1 Absorbance [FeSCN 2+ ] equil [Fe 3+ ] equil [SCN - ] equil K 2 3 Average K How is [FeSCN 2+ ] equil obtained? Show your calculation for Solution 1.

4 22 Chem 1B Dr. White How are [SCN] equil and [Fe 3+ ] equil obtained? Show your calculations for Solution 1. Show how you determined the K value for Solution 1 1: Post Lab Questions 1. Based on your average value of the equilibrium constant for this reaction, classify this reaction as either reactant or product favored and explain your choice. 2. Were the calculated equilibrium concentrations of each of the three chemical species the same in your three equilibrium solutions? Should they have been? Explain. 3. Were the calculated equilibrium constants the same in your three equilibrium solutions? Should they have been? Explain.

5 Chem 1B Dr. White A key part of lab is understanding what effect mistakes may have on results. If you added too much nitric acid to test tube 1, how would this affect the equilibrium constant you determined for solution 1? Explain your answer. (Saying their results would be off or altered is not sufficient. Be specific. For example, would K be larger, smaller, or would there be no change?) 5. A group of students made 5 standard solutions and measured their corresponding absorbance values at a wavelength of nm to generate the graph below. The colorless reactants Y and Z form unknown X from the reaction: 2Y + Z X A solution is prepared by mixing together the following: Volume M Y (ml) Volume M Z (ml) This solution reaches equilibrium, and then is placed in a 1.00 cm wide cuvet and inserted into the spectrometer, producing an absorbance reading of at a wavelength of nm. Calculate K for the above reaction.

6 24 Chem 1B Dr. White

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