EXPERIMENT 3 THE IODINE CLOCK

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1 EXPERIMENT 3 THE IODINE CLOCK Introduction The Rates of Chemical Reactions Broadly defined, chemical kinetics is the study of the rates at which chemical reactions proceed. Oftentimes, reaction rate data helps chemists to develop reaction mechanisms for a given chemical reaction. The rate of a chemical reaction depends upon several parameters, such as temperature, pressure and the initial composition of a system undergoing a reaction. If one wishes to consider a reaction for which a balanced chemical equation is known, several reaction rates can be determined: [1] In [1], it is observed that 2A*s and 4B*s are required to produce 3C*s and one D. For dilute solutions, reaction rates can be expressed by considering the change in concentration of a particular constituent that is being consumed or produced with respect to the change in time: Reaction Rate Law and Reaction Order The rate, ν, at which a chemical reaction proceeds is often found to be proportional in some fashion to the concentration of each reactant raised to some exponential power: [2] Where a and b are experimentally determined values and k is the rate constant. The rate law is not, in general, to be inferred from a balanced chemical reaction, rather, it must be determined experimentally. The constants a and b are oftentimes integers and the overall reaction order is equal to their sum. Thus, for example, if a and b were determined to be equal to 2 and 1, respectively, the reaction is said to be third-order (a + b = = 3). Determination of the Rate Law One very popular method of determining the rate law is known as the method of initial rates, and is aptly titled as the rate is measured only for the initial stages of a chemical reaction. This method involves making a series of solutions, each with a different composition and observing the initial reaction rate as a function of the changing reactant composition. For a reaction which involves two reactive species, the general process involves first keeping the concentration of one of the reactants constant (reactant A, for example) while varying the concentration of the other (reactant B). Then, the concentration of B is held fixed while varying the concentration of A. By plotting the logarithm of the initial rate, ν 0, against the logarithm of the concentration of the reactant whose concentration is being varied, the reaction order with respect to that constituent can be determined as it is equal to the slope of a first-order line of best fit (m = a, according to [2], for the case where [A] is being varied; see Figure 3.1). 29

2 log v 0 (units of M s -1 ) log(v 0 ) = log([A]) R 2 = log [A] (units of M) Figure 3.1 Plot of log ν 0 versus log [A]. According to the line of best fit, the reaction is clearly second-order with respect to reactant A. The Iodine Clock In this lab experiment, the reaction between the persulfate ion (S 2 O 8 2 ) and the iodide ion (I ) is studied in aqueous media and may be represented by the following chemical equation: [3] Two additional reagents are used: the first provides the thiosulfate anion (S 2 O 3 2 ). This anion reacts very quickly, with respect to the process described by [3], with any iodine present: [4] Once all of the thiosulfate is consumed, the iodine being evolved at this point will react with any iodide present and their product forms a complex with the second additional reagent, which is starch. This reaction produces a purple colour which signals the end of the data sampling period. Therefore, by considering the amount of thiosulfate added to the reaction mixture, one can measure the rate of the reaction. 30

3 The Reaction Medium One should also consider the influence that the reaction medium has upon the reaction rate. The observed reaction rate is not only a function of the items outlined earlier, but is also dependent upon the solvent used (not considered in this experiment) and the ionic strength, I, of the solution. In order to determine the rate constant at infinite dilution, k 0, a correction which accounts for the ionic strength of the solution is done (note that this form of the correction is only applicable when the Debye-Hückel limiting law holds): Where z A and z B are the charge numbers associated with the two atoms or molecules that are undergoing the chemical reaction under study. The Rate Constant as a Function of Temperature The rate constant for many chemical reactions is found to increase in a linear fashion when a plot of ln k versus 1/T is composed. This is the graphical expression of the Arrhenius equation: [5] [6] where A is commonly denoted as the frequency factor and where E a is the activation energy of the reaction under consideration. Activated Complex Theory In bimolecular reactions, the two reacting species (A and B, for example) are postulated to pass through some form of transition state, which is often higher in energy than the reactants. This is the basis of activated complex theory (sometimes referred to as transition state theory). By using this theory, one can relate the rate constant to the Gibbs energy of activation, G, for a particular temperature: Where κ is the transmission coefficient, a dimensionless quantity which will be assumed to equal 1, k is the Boltzmann constant, h is the Planck constant, and H and S are the enthalpy of activation and the entropy of activation, respectively. All other parameters carry their usual meaning. For reactions carried out in solution, the enthalpy of activation is related to the activation energy: Therefore, with G and H known, S can be determined by using the equation which expresses the change in Gibbs energy for an isothermal process. [7] [8] 31

4 Materials 4 50 ml burettes 3 medium test tubes 1 test tube rack ml beakers 1 1 L beaker ml volumetric flasks ml Erlenmeyer flasks 1 hot plate & stirrer 1 water bath 1 thermometer 2 scoopula 2 magnetic stir bars 1 stopwatch ml pipette & bulb ml pipettes ammonium chloride (NH 4 Cl) potassium chloride (KCl) ammonium persulfate ((NH 4 ) 2 S 2 O 8 ) potassium iodide (KI) sodium thiosulfate pentahydrate (Na 2 S 2 O 3 C5H 2 O ) starch solution Procedure The following solutions will be provided: M ammonium chloride (NH 4 Cl) (A1) M potassium chloride (KCl) (A2) A. Preparation of Required Solutions Prepare the following solutions using the room temperature de-ionized water (note: DO NOT use tap water for the preparation of these samples): ml of 0.05 M ammonium persulfate ((NH 4 ) 2 S 2 O 8 ) (A3) ml of 0.12 M potassium iodide (KI) (A4) ml of ca M sodium thiosulfate pentahydrate (Na 2 S 2 O 3 C5H 2 O) (A5) 32

5 B. The Effect of Reactant Concentrations on ν 0 1. For A3 and A4, transfer some of each solution into separate 100 ml beakers and then into separate 50 ml burettes. You may wish to label the burettes so you do not mix up the solutions. 2. In a similar fashion as step B1, transfer some de-ionized water, as well as some starch solution, into 50 ml burettes. A total of four burettes should be filled with solution. 3. Prepare two solutions as indicated by the trial #1 row in the table below (do not mix them!): #1 - Into Erlenmeyer Flask #2 - Into Test Tube Trial # A1 A3 A5 H 2 O A2 A4 Starch 1 0.0* * amounts of added solution are to be taken as in units of ml 4. While the solutions for this trial are being prepared, the hot plate can be setup as depicted in Figure 3.2 and the stopwatch prepared for use. 5. Once a stable temperature has been established (do not turn the heating function on), record this temperature in your data sheet, immerse the Erlenmeyer flask into the water bath and when ready, quickly add the contents of the test tube to the flask. Start timing about halfway into adding the contents of the test tube. 6. You will notice that the solution is not coloured in any way, yet appears slightly cloudy. Stop timing when the solution is purple in colour (i.e., continue timing if it is only faintly purple). 7. After the trial, clean and dry the test tube and flask as best as you can, then transfer the items into the oven to eliminate any residual water. 8. Repeat steps as necessary until you have time measurements for all 7 trials. 33

6 test tube stand Erlenmayer flask thermometer stir bar water bath hot plate / stirrer Figure 3.2 C. The Effect of Temperature on ν 0 1. Fill the 1 L beaker about 3/4 of the way with de-ionized water. This beaker will serve as the new water bath for the remainder of this experiment. 2. Heat the bath until the temperature is around 35 EC and then remove the beaker from the heat source. Prepare the solutions using the same composition as outlined in trial #1 and hence, conduct trial #1 once again at this temperature. Record the initial temperature, the elapsed time and the final temperature (the temperature after colour change) on your data sheet. The average of the starting and final temperatures will be used as the temperature at which the trial was conducted. 3. Repeat step C2, but this time heat the bath until the temperature is around 45 EC. Calculations Part B. 1. For each trial in Part B, determine the initial concentration of all the ionic species present and tabulate your results. You might find it useful to explicitly write down all of the chemical dissociation equations. Determine the average reaction temperature. 2. Determine the initial reaction rate, ν 0, for all trials in part B. Express these rates in units of M min Create two plots, one of log ν 0 versus log [I ] using the data from trials 1 4 and one of log ν 0 versus log [S 2 O 8 2- ] using the data from trials 1 and 5 7. Fit the data using a function of the form y = mx + b. From the slope, determine the reaction order with respect to each component. Round your value to the nearest half-integer. 4. Determine the rate constant for each trial and average these values. This mean value will be known as the average rate constant at room temperature, k rt. 5. Determine the ionic strength of one of the solutions used and using equation [5], determine 34

7 the rate constant at infinite dilution, k 0, and compare this value with the literature. Assume that the Debye-Hückel limiting law holds. Part C. 1. Determine the rate constant of the experiment at the two elevated experimental temperatures. Use these two values and k rt to create a plot of ln k versus 1/T and determine E a for this reaction. Interpolate to find k 298 (i.e., the rate constant at T = 298 K). 2. Using k 298 and equations [7] and [8], determine GE and HE and comment upon the chemical significance of these values. Note: pay close attention to units when solving [7] for GE. Use these two values to determine SE and comment upon the physical meaning of the value isolated. 3. Using an appropriate reference (see list below), determine the standard Gibbs energy of reaction, r G, from the f G values associated with each reacting species. 4. Using the values determined in steps C2 and C3, sketch a Gibbs energy reaction profile as the reactants are converted into products. Lab Questions 1. Starting from the form of the rate equation given in [2], show why a plot of log ν 0 vs. log [A] would give the order of reaction with respect to that particular reactant. 2. Why were the NH 4 Cl and KCl salts used? If water were to be substituted for them, how would it affect the observed rate constant and by how much? (find the percent difference) Hint: assume the same starting composition for one of the trials where a non-reagent salt was used, but replace the added salt with water. 3. If reaction [3] is the third-order reaction, what are the units of k(t) [3]? 4. An investigator is conducting kinetic experiments on a first-order reaction A P for which the rate coefficient at 298 K is k 0 = s -1. The activation energy for the reaction is found to be 40 kj mol -1. Assume that a simple Arrhenius expression gives the dependence of k upon temperature. The investigator runs the reaction in a container whose temperature is being continuously varied such that the temperature at time t is T = (T 0 )/( t) where T 0 is the temperature at time t = 0. When half of the initial amount of A has been consumed, what is the temperature in the reaction vessel? References 1. Atkins, Peter and Julio de Paula. Physical Chemistry. 7 th ed. New York: W. H. Freeman, , , , , Bernasconi, Claude F., ed. Investigation of Rates and Mechanisms of Reactions. 4 th ed. Toronto: John Wiley & Sons, Wagman, Donald D. Selected Values of Chemical Thermodynamic Properties. Washington: National Bureau of Standards, Howells, W. J. Journal of the Chemical Society (Resumed). 1939, Amis, Edward S., and James E. Potts. J. Am Chem. Soc. 1941, 63, Moews, P.C., and R. H. Petrucci. Journal of Chemical Education. 1964, 41,

8 Data Sheet A. Concentration of NH 4 Cl solution: mol L -1 Concentration of KCl solution: mol L -1 Concentration of (NH 4 ) 2 S 2 O 8 solution: mol L -1 Concentration of KI solution: mol L -1 Concentration of Na 2 S 2 O 3 solution: mol L -1 B. Trial # Temperature (EC) Elapsed Time (m:s) Average Reaction Temperature: EC C. Trial T i (EC) T f (EC) T& (EC) Elapsed Time (m:s) ca. 35 EC ca. 45 EC 36

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