CHEM 3364 Spring 2008 1 Kinetics of an Iodine Clock Reaction Objective: Determine the experimental rate law for a peroxysulfate decomposition reaction. Find its activation energy. Consider appropriate reaction mechanisms that match the experimental rate law. Introduction: Thermodynamics allows us to predict whether a reaction is favorable. Exactly how the reaction occurs on a microscopic level, however, cannot be determined from ∆G. The process by which a reaction occurs, what we call the reaction’s mechanism, can often be deduced by measuring how fast the reaction occurs. We can do this because a reaction’s rate depends on the rate of effective collisions between reacting species. This collision rate is directly related to the concentrations of those species. The overall reaction rate is obtain by appropriately combined the collision rates of species that participate in any step up to and including the slowest or rate-determining step. The relationship between the reaction’s rate and the concentrations of species affecting the rate, R, is given by the rate law, which takes the general form R = k[A]α •••[D]δ where k is the reaction’s rate constant and the terms in brackets, [ ], are concentrations. The superscripts α…δ are called reaction orders. The purpose of this experiment is to characterize the reaction between iodide, I-, peroxydisulfate (or persulfate), S2O82-, and thiosulfate, S2O32-. This reaction system is an example of what is commonly known as a “clock reaction.” At one step during the reaction, a product is made which reacts with an indicator solution; the resulting change in color is used to time the reaction. Preliminary work will focus on identifying the individual steps comprising the overall reaction. An investigation of the reaction’s kinetics provides additional information. Solution Preparation You will need to prepare 200 mL each of the solutions listed below. With the exception of the peroxysulfate solution that must be prepared each week, these solutions, if well sealed, can be stored and used for several weeks. If you run out of any solution, you can prepare an additional amount. Be sure to note, however, the cautionary statements in Tasks I and II. 0.20 M KI 0.20 M KCl 0.0050 M Na2S2O3 0.10 M Na2SO4 0.10 M K2S2O8 (this must be prepared each week) Prepare these solutions using appropriate volumetric glassware. Two additional solutions – a saturated solution of I2 and a starch indicator solution – are available in the lab.CHEM 3364 Spring 2008 2 Preliminary Investigations Much of the chemistry in this system can be deduced using simple experiments in which two or more reagents are mixed together and the results recorded. Each preliminary investigation is followed by one or more questions. The answers to these questions are important so give them careful consideration. When asked to write tentative reactions, put them in the following general form: R1 + R2 → P1 + P2 where R represents a reactant and P represents a product. Whenever possible, replace R and P with known reactants and products. If you cannot identify a reactant or product (but know that one must be there) then retain the needed R and/or P designations in your reaction. If possible, balance your reactions. Investigation 1. Place approximately 5 mL each of the KI, KCl, K2S2O8, Na2S2O3, and I2 solutions in separate test-tube. To each test-tube, add a few drops of the starch indicator solution and record your observations. Questions: What reagent is responsible for changing the indicator’s color? Is this reagent a reactant or a product? Write a tentative reaction that explains your observations using starchI to represent the indicator’s colorless form and starchC to indicate its colored form. We will assume that the starch indicator always reacts in a 1:1 stoichiometry. Investigation 2. Mix together ten different combinations (two-at-a-time, about 5 mL each) of your reagents. Add a few drops of the indicator solution to each combination and record your observations. Questions: Which combinations of reagents produce a change in the indicator’s color? Write a tentative reaction that explains your observations. Investigation 3. In separate test tubes place (a) about 5 mL of KI, (b) about 5 mL of KI and 1 mL of Na2S2O3, and (c) 5 mL of KCl. Add a few drops of the indicator solution to each test tube. Finally, add about 5 mL of K2S2O8 to each test tube and record your observations. As an additional experiment, try adding 1 mL of Na2S2O3 to each test tube and recording your observations. Questions: Which combinations of reagents produce a change in the indicator’s color? For those combinations that provided a positive result, are there ways in which they behaved differently? How can you explain this difference? Write tentative reactions that explain your observations. What have we learned to this point? What is the chemical reaction with the K2S2O8? How does Na2S2O3 participate? What is its role?CHEM 3364 Spring 2008 3 Measuring the Reaction’s Rate – Task I To study a reaction kinetically we must determine the reaction’s rate. One approach is to measure the time required to react a known amount of reactant or the time required to make a known amount of product. This, of course, gives us an average rate,ν, over the elapsed time [concentration]t∆ν=∆ where ∆[concentration] is the change in concentration and ∆t is the change in time. The color change you observed in Task I can be used for this purpose. Use the following procedure as a guide. Use a thermistor-controlled water bath to maintain a constant temperature approximately at room temperature. Using appropriate volumetric glassware, transfer 20.0 mL of 0.20 M KI, 10.0 mL of 0.0050 M Na2S2O3, and 3-4 drops of indicator into an Erlenmeyer flask. Into another Erlenmeyer flask, add 20.0 mL of 0.10 M K2S2O8. Put both flasks in the water bath and allow them to come to thermal equilibrium. Mix the two solutions together and begin timing. While swirling the flask to ensure that the solutions are well mixed, observe the solution and note the time required for the indicator to change color. Repeat the procedure several times until you are confident that you can measure ∆t reproducibly. Be sure to calculate your average value for ∆t. Now, calculate the