A Chemical Clock Things to Consider 1. What are the three major objectives of this experiment? What methods will you try using to achieve each of these three objectives? 2. What is difference between reaction time and reaction rate? What are the units of each? How will you convert the reaction times you measure into reaction rates? 3. What is the role of the thiosulfate ion in the Landolt Clock? 4. When setting up your reaction mixtures for the Landolt Clock, which solution concentrations must be varied and which must be held constant? Why? 5. Consider each of the following questions regarding data and measurements: a) When you are making a measurement that will be used directly in a calculation, what is the minimum number of trials should you conduct in order to ensure reasonable precision and accuracy of your results? Why? b) What is the minimum number of data points needed to construct a plot where the equation of the best-fit-line has reasonable precision and accuracy? Why? c) Do you need to conduct multiple trials when making measurements that will be plotted as a graph? Why or why not? 6. Referring to Table 1, Starting Mixtures in Your Team s Project, why is water added to Mixtures A and C, but not to B? 7. Create a Table of Mixtures similar to Table 1, Starting Mixtures in Your Team s Project detailing the actual mixtures you will use with the Method of Initial Rates to determine the experimental rate law. Be sure to consider your answers to Questions 3, 4, and 5 above when constructing your team s table. 8. Express Equation (4) as a linear equation in! = #$ + & format. How can you use this linearized expression to find the value of the activation energy, E a, for your reaction? 9. Do any of the chemicals or equipment used in this experiment require special handling or disposal? If so list each and what special precautions are indicated. A Chemical Clock Your Team s Project Page 1
A Chemical Clock Your Team s Project Your team of chemical consultants has been hired by, I-now a new biotech firm that is planning to use a variation of the Landolt Clock to produce molecular iodine at a controlled rate from a medical device they are developing. The company needs your team to determine the chemical rate law, rate constant, and the dependence of the rate constant on temperature for the reaction so that they can optimize the reaction kinetics for use in their device. They would also like you to evaluate the effectiveness of a new ammonium molybdate catalyst on the reaction rate. You will have two weeks to complete your project. The chemical equation describing the reaction used in their device is, 6 I +, + BrO 0 1 +, + 6 H 3 +, 3 I 6 +, + Br +, + 3 H 6 O(8) (1) This reaction can be coupled with a second one described by the equation, I 6 +, + 2 S 6 O 0 61 +, 2 I 1 +, + S < O = 61 (+,) (2) and a starch indicator to form a variant of the Landolt Clock. In this dramatic reaction discovered by Swiss chemist Hans Heinrich Landolt in 1886, the reactants are mixed to form a colorless solution that after a measureable time delay suddenly turns dark blue due to the formation of a starch-iodine complex. Notice that molecular iodine, I 2, is produced by the reaction described by Equation (1), and then removed by the reaction with thiosulfate, S 2 O 3 2-, described by Equation (2). The quantity of thiosulfate added acts as a timer that determines how long it takes before the dark blue starch-iodine complex forms. This process works because the rate of the reaction described by Equation (2) is much faster than that described by Equation (1), so that the iodine produced by Equation (1) is removed from the solution by the thiosulfate present through Equation (2) before the iodine can complex with the starch indicator and turn blue. This continues until all the thiosulfate added has reacted, at which point the iodine formed in Equation (1) is no longer removed from solution and so complexes with the starch indicator turning the solution blue. Thus, the time interval between mixing the two solutions and the appearance of the blue complex depends upon the rate of the reaction and the amount of thiosulfate added. By keeping the amount of thiosulfate constant in each trail, we can use the time it takes for the mixture to turn blue as a measure of the relative rate of the reaction under various conditions. As a simple analogy of how the thiosulfate in the Landolt Clock works, consider a leaky faucet for which you wish to measure the leak rate over a short period of time. One method would be to place a small bucket of known volume below the faucet and measure the time that it takes to fill this bucket. The thiosulfate in the Landolt Clock takes the part A Chemical Clock Your Team s Project Page 1
of the bucket. The iodine produced in Equation (1) drips into the reaction described by Equation (2) until the thiosulfate present reacts with as much iodine as it can handle and the bucket becomes full, as signaled by the blue color change. By measuring the time it takes for a set concentration of thiosulfate ions to be consumed, we can measure the relative rate of Equation (1). Assuming the reaction described by Equation (1) is under initial rate conditions when the color change occurs, we can systematically vary the concentration of each of the three reactants to determine the reaction rate law using The Method of Initial Rates described in your textbook. The reaction rate law for Equation (1) can be expressed as: >+?@ = A I 1 B BrO 0 1 C H 3 D (3) where x, y, and z are the orders of the reaction with respect to the iodide, bromate, and hydrogen ion concentrations, and k is the rate constant. Your team will first need to design a set of experiments to determine the values of each of these four values at a given temperature. Once your team has determined the reaction rate law, you will then need to determine the value of the activation energy of the reaction in order to find the dependence of the value of the rate constant, k, on temperature. The rate constant can be expressed by the Arrhenius Equation as, A = E@ 1F G/IJ (4) where T is the temperature, E a is the activation energy, A is the pre-exponential factor, and R is the universal gas constant. By determining the activation energy of the reaction, you can predict the value of k at a given temperature, and therefore the rate of the reaction as described by Equation (3) under a specific set of conditions. Finally, it will be up to your team to determine the effects of an ammonium molybdate catalyst on the rate law and the activation energy. You will also need to research the safety of this catalyst for use in biological systems. A Chemical Clock Your Team s Project Page 2
Getting Started Your team will have the following chemicals available to work with: 0.0010-M Sodium thiosulfate, Na 2 S 2 O 3 (aq) 0.010-M Potassium iodide, KI(aq) 0.040-M Potassium bromate, KBrO 3 (aq) 0.10-M Hydrochloric acid, HCl(aq) 0.5-M Ammonium molybdate, (NH 4 ) 2 MoO 4 (aq) 3% Starch indicator Safety Notes: Potassium bromate is a potential carcinogen and gloves should be worn when handling this. Ammonium Molybdate is an environmental toxin and must be disposed of properly. Wear chemical safety goggles at all times. Dispose of all chemical waste properly and follow all safety protocols outlined by your laboratory institution and instructor. To see how the Landolt Clock works, obtain two clean dry 125-mL Erlenmeyer flasks. Label these Flask I and Flask II. Prepare Mixture A listed in Table 1 below as follows: Into Flask I combine 10-mL of 0.010-M potassium iodide, 10- ml of 0.0010-M sodium thiosulfate, and 10-mL of deionized water. Into Flask II combine 10- ml of 0.040-M potassium bromate, 10-mL of 0.10-M hydrochloric acid, and 3-4 drops of the 3% starch indicator. Measure the temperature of the mixtures in Flask I and Flask II either using two separate thermometers or being careful to rinse a single thermometer between measurements in order to avoid contamination of the mixtures in either flask. Both flasks should be at the same temperature. If not, adjust the temperature of the flasks using warm or ice water baths so that they are the same. Record this value as the initial temperature. Leave the thermometer in Flask I so that you can measure the final temperature at the end of the reaction. Table 1: Starting Mixtures Flask I (125 ml) Flask II (125 ml) 0.0010-M 0.010-M 0.040-M 0.10-M 3% Mixture: Na 2 S 2 O 3 KI H 2 O KBrO 3 HCl Starch A 10.0 ml 10.0 ml 10.0 ml 10.0 ml 10.0 ml 3-4 drops B 10.0 ml 20.0 ml 0 ml 10.0 ml 10.0 ml 3-4 drops C 10.0 ml 5.0 ml 15.0 ml 10.0 ml 10.0 ml 3-4 drops A Chemical Clock Your Team s Project Page 3
Your team will need a stopwatch, timer, or clock with a second hand to measure how long it takes for the reaction to turn blue after mixing. Do your best to start your timer the moment the two solutions are mixed and to stop your timer the moment the blue color first appears. Starting your timer, pour the contents of Flask II into those of Flask I. Gently swirl the contents to mix. Mixture A should take about a minute or so to change color at room temperature. Record the elapsed time. Measure the final temperature of the mixture after the color change occurs and record this as the final temperature. If the initial and final temperatures differed slightly you can simply average the two temperatures and record this as the reaction temperature; if they differed by more than a few degrees repeat the trial keeping the temperature of the mixture as constant as possible by immersing the reaction flask in a hot or cold water bath during the reaction as needed. The concentration of iodide in Mixture A is given by: I 1 = KL.L ml L.LKL mol LTU VL.L ml = 0.0020 mol L 1K ( Mixture A) (5) Similarly, the concentration of iodide in Mixtures B and C are: I 1 = 6L.L ml L.LKL mol LTU VL.L ml = 0.0040 mol L 1K ( Mixture B) (6) I 1 = V.L ml L.LKL mol LTU VL.L ml = 0.0010 mol L 1K ( Mixture C) (7) Mixtures B and C are given here as examples of mixtures your team can use to determine the reaction orders using the method of initial rates. Notice that by doubling or halving the volume of potassium iodide in these mixtures compared to that in Mixture A, we double or halve the iodide concentration in the final reaction mixture, respectively. This is achieved by adjusting the volume of water so that the total volume of the final reaction mixture remains constant. Proceeding in this way, your team can investigate the reaction kinetics of the uncatalyzed reaction and the effects of temperature on the reaction rate. When your team is ready to analyze the effects of the catalyst begin by creating Mixture A again. Add three drops of the 0.5-M Ammonium Molybdate catalyst to the mixture in Flask II and observe the effects of the catalyst on the reaction rate when the contents of the two flasks are combined. A Chemical Clock Your Team s Project Page 4
A Chemical Clock Your Report: Part I: Determination of the Rate Law 1. Complete Table 1 below using the volumes of each chemical used and the reaction times measured that your team used to determine the rate law. Leave any unused rows blank. Table 1: Mixtures used to Determine the Rate Law Mixture: 0.0010-M Na 2 S 2 O 3 Flask I 0.010-M KI H 2 O 0.040-M KBrO 3 Flask II 0.10-M HCl Starch (drops) 1 10.0 ml 10.0 ml 10.0 ml 10.0 ml 10.0 ml 3-4 2 3 4 5 6 7 8 9 10 11 12 Temp / C Time / s 2. Using the data in Table 1, calculate the initial concentration of thiosulfate ion, [S 2 O 3 2- ] 0 in Mixture 1 (show all work and units): 3. Using the concentration of thiosulfate ion in Mixture 1 and your measured reaction time, determine the initial reaction rate for Mixture 1 (show all work and units): A Chemical Clock - Your Report Page 1
4. Complete Table 2 below by determining the concentration of each ion in solution and the corresponding reaction rate for each mixture. Leave any unused rows blank. Mixture 1 2 3 4 5 6 7 8 9 10 11 12 Table 2: Initial Concentrations and Rates [S 2 O 3 2 ] 0 /M [I ] 0 / M [BrO 3 ] 0 / M [H + ] 0 / M Rate / M s -1 Temp / C 5. Use the method of initial rates and the relevant data from Table 2 above to determine the order of each reactant and the experimentally determined rate law. Attach one or more sheets neatly detailing your work. Be certain to include units in your calculations. The Rate Law: A Chemical Clock - Your Report Page 2
6. Using your experimentally determined rate law and the data in Table 2, determine the value and units of the rate constant, k, for Mixture 1. Show your work. 7. Complete Table 3 below by determining the value of the rate constant, k, for each of your reaction mixtures and the average value of the rate constant for these trials. Leave any unused cells blank. Record the average temperature at which these values were measured. Table 3: Value of the Rate Constant Mixture Value of k Mixture 1 7 2 8 3 9 4 10 5 11 6 12 Value of k Average value of k: Units of k: Average temp: 8. The value of the rate constant, k, for each of these reaction mixtures should be similar. Briefly explain why. A Chemical Clock - Your Report Page 3
Part II: Determination of the Activation Energy, E a 9. Set up your experiments and make measurements to determine the value of the activation energy, E a, of the reaction. In Table 4 list the experimental conditions and data your team measured. In Table 5 list the results calculated from these data and the value of the activation energy you determined for each mixture. Fill in the appropriate headings for each column and indicate all units. Leave any unused cells blank. Mixture 1 Table 4: Data Measured for Determining the Activation Energy 2 3 4 5 6 7 8 Mixture 1 Table 5: Results used for Determining the Activation Energy 2 3 4 5 6 7 8 Average Activation Energy: Units: 10. Attach one or more additional pages detailing how you used the data and results in Tables 4 and 5 to determine the activation energy above. Include a brief explanation of your method. Include sample calculations, graphs, or other supporting work. A Chemical Clock - Your Report Page 4
Part III: The Catalyst Using one or more sheets, attach answers to the following questions regarding your team s investigation of the effects of the ammonium molybdate catalyst on the reaction rate law: 11. Briefly describe the specific challenges you encountered in determining: (a) the rate law, and (b) the activation energy, for the catalyzed reaction and how you overcame these. 12. Attach your own neatly organized data and results tables for your team s study of: (a) the rate law, and (b) the activation energy, for the catalyzed reaction. Include a brief explanation of your methods. Include sample calculations, graphs, or other supporting work. Be sure to include your final experimentally determined rate law for the catalyzed reaction and the value you determined for its activation energy (and units). 13. Based on your team s results, is the role of ammonium molybdate in the Landolt clock reaction consistent with the definition of a catalyst? Explain why or why not. 14. Ammonium molybdate is described in this experiment as, an environmental toxin, but is it safe for use inside I-now s proposed medical device at the low concentrations used in this experiment? Look up the safety of ammonium molybdate in biological systems and give your teams evaluation in a brief one-paragraph summary. Be sure to properly cite your sources and support your assertions using quotes and any needed calculations. Be as quantitative as possible in your evaluation. A Chemical Clock - Your Report Page 5