Energy and Equilibrium in Chemical Systems Iodine Clock Kinetics Lab Introduction Chemical kinetics is the study of the rates of reactions. In this experiment, we will investigate the effect of changing certain conditions of a reaction on the rate of the reaction. The reaction we are using is called the Harcourt-Esson reaction, or more commonly the Iodine Clock reaction. The reaction proceeds in two steps: Step I (Production of triiodide ion) H2O2 (aq) + 3 I- (aq) + 2 H+ (aq) I3- (aq) + 2 H2O (l) Step II (Restoration of iodide ion) I3- (aq) + 2 S2O3-2 (aq) 3 I- (aq) + S4O6-2 (aq) We can visualize the progress of the reaction by introducing an indicator, starch, which turns a deep blue or purple in the presence of the triiodide ion. Therefore, the reaction starts clear (thiosulfate is reacting with triiodide to reduce it to iodine) and become purple (thiosulfate has been exhausted). We will investigate the effects of several changes in solution composition as well as the addition of energy as heat on this reaction. Purpose Students will design and execute experiments to determine what factors affect the rate of the iodine clock reaction. Materials 0.05 M KI 1% starch 0.01 M Na2S2O3 1.0 M H2SO4 0.05 M H2O2 Stopwatch Hot plate Ice bath 100 mL (or larger) Erlenmeyer flask Procedure “Base” reaction 15 mL KI solution 5 mL starch solution 2.5 mL sodium thiosulfate solution 5 mL sulfuric acid solution 15 mL distilled water 15 mL hydrogen peroxide Total solution volume: 57.5 mL Part I Using the materials of the lab, write your own procedure to determine what affects the rate of reaction, and to quantify the affect you observe. Focus on the concepts underlying a good scientific experiment – repeatability, thoroughness, control of variables. Check your procedure with the teacher before performing it. Add the substances in the order listed to a clean Erlenmeyer flask. Begin timing the reaction once the hydrogen peroxide first touches the solution. Stop timing the reaction when a visible color change occurs. Repeat three times to establish a good baseline. Part II Perform the procedure you wrote during Part I. Gather data to determine how the reaction rate is changed over time. Lab Analysis Questions 1. Calculate the concentration of each substance in the “base” reaction. Starch concentration may be calculated as a mass percent. 2. Quantify the effect each change you made had on the rate. For example, “When the concentration of sodium thiosulfate was increased by a factor of 2, the rate decreased by a factor of 4.” 3. Estimate the order of reaction with respect to each species that had an effect on the reaction rate and propose a rate law for the reaction. 4. A second proposed mechanism for the reaction is shown below. Which mechanism, the one below or the one at the beginning of this document, is most likely to be correct according to your data? Defend your answer in the context of your data and the principles of chemical kinetics. Step 1: H2O2 + I- H2O + IOStep 2: H+ + IO- HIO Step 3: H+ + I- + HIO I2 + H2O (slow) Teacher’s Notes - One key concept here that keen students may pick out is that the thiosulfate ion is the “fuse” for this reaction. The blue color cannot appear until the thiosulfate ion is used up. This concept is critical, and it helps to take steps to ensure each student understands it. In fact, it would be quite interesting to gently nudge one group into performing the reaction without the thiosulfate. - So long as the thiosulfate ion concentration is kept constant, the kinetics of the first reaction can be investigated. Changing the thiosulfate ion concentration changes the “delay”, which may seem to be affecting the first reaction but in reality is not.