Iodine Clock Reaction Introduction In this experiment you will determine the Rate Law for the following oxidation reduction reaction 2 H aq 2 I aq H2O2 aq I 2 aq 2 H2O l 1 The rate or speed of the reaction is dependent on the concentrations of iodide ion I and hydrogen peroxide H2O2 The spectator ions are left off the reaction Therefore we can write the Rate Law concentration dependence for the reaction as Rate k I x H2O2 y 2 Where x is the order of the reaction in I y is the order of the reaction in H2O2 and k is the rate constant The temperature dependence of the rate is seen in k that is there is a separate value of k for each temperature at which the reaction takes place The temperature must therefore be held constant to accurately calculate x y and k Since the Rate Law is empirical we have to go to the lab to make measurements that will enable these values to be calculated The rate will be measured for the reaction near time 0 so that few products been formed and there will be no reverse reaction The concentrations of iodide and hydrogen peroxide will be varied and the rates compared to find each order i e the exponents x and y This is the Method of Initial Rates and it will be used to find x y and k As with a lot of kinetics the concentration of reactants or products at any instant is difficult to measure directly so in this lab the rate will be determined indirectly We have a very handy test for the presence of one of the products iodine I2 namely starch Starch reacts with iodine to form a blue black colored complex Unfortunately as soon as any iodine is produced it will react to make the complex and the solution will turn blue black instantaneously Thus using starch as an indicator by itself would not be of much help It confirms that some amount of I2 is being formed but it tells us nothing about what we are trying to measure the rate how long it takes to produce a given quantity of I2 To get around this problem we will introduce a side reaction that will remove the initial I2 that is produced by our main reaction This will prevent the solution from turning black for long enough so that we can make some good time measurements We will use the following side reaction I2 aq 2 S2O3 2 aq 2 I aq S4O6 2 aq 3 S2O3 2 thiosulfate ion reacts with I2 which prevents the solution from turning blue black How will this help Since we have carefully measured the amount of the thiosulfate a small amount that will run out fairly quickly we know exactly how much iodine it will take to react with this thiosulfate As soon as all of the thiosulfate has reacted iodine will then be left over and then the reaction will change color By putting in this time delay we can now calculate the rate at which I2 is being formed The length of time required for the reaction to change color will be the time that you will measure and the amount of I2 produced in that time is the amount of I2 that is needed to react with the thiosulfate So now we have the rate of formation of I2 and depletion of thiosulfate and can calculate the final rate value we are looking for in this lab which is the moles of hydrogen peroxide reacted per second Rate moles H2O2 reacted time in seconds 4 1 mole H2O2 reacts to produce 1 mole I2 which reacts with 2 moles S2O3 2 See reactions 1 and 3 So 1 mole H2O2 2 mole S2O3 2 Since the moles of S2O3 2 in solution is known we can calculate it we know that the number of moles of H2O2 that react will be of the moles of S2O3 2 in the solution The rate is then this number of moles of H2O2 divided by the time Calculate the moles of S2O3 2 from the volume and molarity of Na2S2O3 2 5 mL of 0 010 M Don t forget that molarity is moles Liter Use this to calcuate the moles of H2O2 reacted Include this number in the final data sheet for Parts I and II Equipment Three 125 or 250 mL Erlenmeyer flasks 10 ml and 5 ml Pipettes Stop watch cell phone Three 100 or 150 mL beakers Thermometer Hot water bath for Part II One bin of chemicals per group that will contain 0 050 M KI 0 050 M KCl 1 0 M H2SO4 Spill B1 1 starch solution 0 010 M Na2S2O3 Ice bath for Part II 0 050 M H2O2 Disposal All mixtures Spill Disposal B1 down the sink Procedure Part I Effect of concentration 1 Clean and mostly dry three Erlenmeyer flasks 2 Label them 1 2 and 3 3 Use fresh pipettes for each solution Rinse pipette twice with the solution that you will be measuring and keep this prepared pipette with the corresponding solution Add the following amounts of the solutions below to prepare each flask The chemicals must be added in the order listed Left to Right Flask 1 2 3 0 050 M KI 15 0 mL 15 0 mL 7 5 mL 1 Starch 5 0 mL 5 0 mL 5 0 mL 0 010 M Na2S2O3 2 5 mL 2 5 mL 2 5 mL 1 M H2SO4 5 0 mL 5 0 mL 5 0 mL 0 050 M KCl 0 0 7 5 ml 4 Rinse and mostly dry 3 beakers and label as 1 2 and 3 5 Prepare the following solutions in clean beakers Beaker 1 2 3 0 050 M H2O2 15 0 mL 7 5 mL 15 0 mL Deionized H2O 0 7 5 mL 0 6 Add the contents of beaker 1 to flask 1 7 Start the stopwatch as soon as you mix the solutions Swirl the flask to mix and note the time it takes for the color to change This is Run 1 Record the temperature 8 Add the contents of beaker 2 to flask 2 Repeat procedure in step 7 This is Run 2 9 Add the contents of beaker 3 to flask 3 Repeat procedure in step 7 This is Run 3 Make sure that the shade of blue black color you observe is the same in each of the three runs 10 The KCl solution and deionized water are added so that the ionic strength and final volume of the various solutions 42 5 mL is kept constant in each run 11 Create a table so that the method of initial rates can be used The initial concentrations of iodide and hydrogen peroxide are not 0 050 M We have to remember that when the solutions were mixed they were diluted Thus Volume 1 Molarity 1 Run 1 Flask 1 beaker 1 Volume 2 Molarity 2 I initial 15 0 mL 0 050M 42 5 ml M2 …