CHM 152LL: Kinetics of an Iodine Clock Reaction Prelab -

Prepare the purpose, brief procedural outline, and data and results tables in your laboratory notebook BEFORE you attend your lab class. Include the information from Table I, the reagents and drops added for Runs 1-5, in your procedure section.

Introduction -

In this experiment, you will determine the rate law for a reaction and the effect of concentration on the rate of the reaction by studying the initial reaction rate at several different reactant concentrations. You will also examine the effect of a catalyst on the reaction rate. Lastly, you will investigate the effect of temperature on the rate of this reaction, which will allow you to calculate the activation energy.

The Clock Reaction and Reaction Rate The primary reaction to be studied is the oxidation of I- (iodide ion) by S2O82- (persulfate ion) in aqueous solution: 2 I-(aq) + S2O82-(aq)  I2(aq) + 2 SO42-(aq)

Equation 1

To find the rate of this reaction, you will need to monitor the change in the concentration of a reactant or product with time. We choose the iodine to monitor because it is a brown color which can be seen. You will do this by using a second reaction, referred to as a "clock" reaction. The clock reaction indicates when a specific amount of I2 has been produced by the primary reaction. If you know how much I2 has formed and the time it takes to produce it, you can calculate the rate of the primary reaction for I2, as shown in Equation 2: Rate =

[I 2 ] t

Equation 2

In Equation 2, [I2] is the change in the concentration of I2 and t represents the corresponding change in time. By examining the initial rates data, you will also be able to determine the rate law for the reaction, which is shown by Equation 3: Rate = k[I-]x [S2O82-] y

Equation 3

In Equation 3, k is the rate constant, and x and y represent the “order” of I- and S2O82- ions, respectively. These values will be determined experimentally from your data. The "clock" reaction you will use involves the reaction of a very small amount of S2O32- (thiosulfate ion) with the I2 produced in the primary reaction: I2(aq) + 2 S2O32-(aq)  2 I-(aq) + S4O62-(aq)

Equation 4 (very fast)

S2O32- is essentially consuming the I2 formed in the primary reaction. S2O32- reacts very rapidly with I2, so I2 is consumed by S2O32- as fast as it is formed. As soon as all of the S2O32- ions have reacted (they are all used up), the I2 still being formed in Equation 1 starts to accumulate and you can finally see the color change. Because small amounts of iodine are faint yellow and hard to see, we will use starch as an indicator to help us “see” the I2 since the interaction between starch and I2 causes a dark blue color. Thus, "∆t" is simply the time elapsed between mixing the reagents and the appearance of the blue color. Because you know the concentration of the S2O32- ions in the reaction mixture, you can calculate "∆[I2]" by using the stoichiometry of the clock reaction. Since the same amount of S2O32- is added to each run, ∆[I2] is also the same for each run. However, the amount of time for the appearance of the blue color varies, so t is not constant. GCC CHM 152LL: Kinetics of an Iodine Clock Reaction

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Note: When the solution turns blue, only a very small percentage of the reactants have been used up. Thus, the initial concentrations of reactants can be used in the rate law with very little error. Typically, the error is less than 5%.

Effect of Concentration on the Reaction Rate: Finding the Rate Law For Runs 1-3, you will vary the initial concentration of I- while keeping the initial concentration of S2O82- constant. By using the method of initial rates, you will be able to determine the order for I- . The following example will illustrate how to find a reaction order using the method of initial rates. Example. Data was obtained for the reaction: A + B  C Experiment 1 2 3

[A], M 0.020 0.030 0.030

[B], M 0.10 0.10 0.25

Rate (M/s) 1.20 1.80 11.25

The general rate law for this example is Rate = k[A]x[B]y Since [A] changes between Experiment 1 and 2, while [B] remains constant, the order for A is obtained by taking the ratio of the rates from these two experiments 1.80 M/s k[0.030] x [0.10] y  1.20 M/s k[0.020] x [0.10] y

Since k is constant at a given temperature and [B]y is constant for Experiments 1 and 2, the equation simplifies to 1.80 M/s  [0.030]    1.20 M/s  [0.020] 

x

or

1.50 = 1.5x

Thus, x = 1 for this example. Unfortunately, experimental results are not usually that "clean", and a more sophisticated method is needed to find x. Mathematically, solving for exponents requires the use of logarithms. Taking the log of both sides of the equation above: log1.50  x  log1.5 

Rearranging this equation to solve for x yields x=

log1.50  =1 log1.5 

Experiments 2 and 3 may then be used to find the order for B, as shown below 11.25 M/s k[0.030] x [0.25] y  1.80 M/s k[0.030] x [0.10] y

By cancelling out the common terms and dividing the rate and concentration values, we obtain 6.25 = 2.5y Taking the log of both sides and rearranging to solve for y gives y=

log6.25  =2 log2.5 

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For Runs 2, 4 and 5, you will vary the initial concentration of S2O82- while keeping the initial concentration of I- constant. By using the initial rates method shown above, you will also be able to find the order for S2O82- . Once you determine the values for x and y, you can substitute these back into Equation 3, and calculate k (the rate constant, with units) for each set of initial concentrations. You will then determine the average value of k for the five room temperature runs.

Effect of a Catalyst on the Reaction Rate A comparison of the reaction rate with and without a catalyst will demonstrate catalytic action for runs 6-9. Cu2+ ions, added in the form of a dilute Cu(NO3)2 solution, makes a suitable catalyst.

Effect of Temperature on the Reaction Rate: Determination of Ea The Arrhenius equation describes the relationship between k, T, Ea and A, the frequency factor: ln k = -

Ea  1    + ln A R T

Equation 6

You will run the reaction at four different temperatures, calculate the rate constant at each temperature, and use an Arrhenius graph of (ln k ) vs. (1 / T) to determine the value of the activation energy, Ea.

Procedure Effect of Concentration – Runs 1-5 Work in groups of 3. You will carry out five runs three times each to give a good average time per run. To keep the ionic strength and volume of all runs relatively constant, the non-reactive compounds KNO3 and (NH4)2SO4 are added to replace the reactants KI and (NH4)2S2O8 when less than 20 drops of either reactant is used. Waste handling: Keep a large beaker at your bench to collect the clock reaction waste. When you have completed all trials, pour the contents of this beaker into the waste container in the fume hood. Table I gives the composition in drops of the reaction mixtures that will be used to determine the order of reaction with respect to the iodide and persulfate ions. Table I. Compostion of the Reaction Mixtures Run No.

0.2% starch

0.012 M Na2S2O3

0.20 M KI

0.20 M KNO3

1 2 3 4 5

3 3 3 3 3

5 5 5 5 5

20 10 5 10 10

0 10 15 10 10

10 mL Beaker

0.20 M 0.20 M (NH4)2S2O8 (NH4) 2SO4 10 10 10 20 5

10 10 10 0 15

Test Tube

Use the dropper bottles containing starch, Na2S2O3, KI and KNO3 to add the appropriate number of drops for Run 1 into one of the 10 mL beakers. Repeat this process with the other 10 mL beakers to prepare the other two trials for Run 1. Note: Make sure you use solutions from the same set of dropper bottles and hold the bottles at the same angle to obtain drops of consistent size! GCC CHM 152LL: Kinetics of an Iodine Clock Reaction

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Use your digital thermometer to record the temperature of the mixture in one of these beakers. Note that this will be room temperature and this temperature will remain relatively constant for Runs 1-5, so only one measurement is needed. Obtain a stirring plate from one of the benchs near the fumehoods. Add a magnetic stirring bar to one of the 10 mL beakers, place the beaker on the magnetic stirrer, and adjust the speed of the stirrer to obtain a slow but steady speed. Use the dropper bottles containing (NH4)2S2O8 and (NH4)2SO4 to combine the appropriate number of drops for Run 1 in a small clean cry test tube. You can use the same test tube for all trials of runs 1-3 because the contents are the same. But you will need a new clean dry test tube for run 4 and another for run 5. Pour the (NH4)2S2O8 and (NH4)2SO4 mixture from the test tube into the 10 mL beaker as rapidly as possible, starting the timer as soon as the contents have been added. Stop the timer when the solution turns blue, and record the time in your data table. Use crucible tongs to remove the stir bar from the beaker to keep your fingers from getting blue. Repeat this process for the other two Run 1 trials, so you will have three reaction times for Run 1. Rinse the beakers and the test tubes with tap water and then deionized water. Dry the beakers. Repeat the procedure for Runs 2-5.

Effect of a Catalyst – Run 6 Repeat Run 3 (at room temperature) one more time, adding one drop of 0.020 M Cu(NO3)2 to the mixture in the test tube. Since there are no calculations using this result, one trial should be sufficient. We will call this run 6.

Effect of Temperature and the Activation Energy – Runs 7-9 You have reaction times for Run 3 at room temperature, and need reaction times for Run 3 at three other temperatures, one at least 10-15°C colder (run 7) than room temperature, one about 10-15°C warmer (run 8) than room temperature, and another about 20-30°C warmer (run 9) than room temperature. Make sure to do the cold temperature run first and only perform one trial at each temperature. The general procedure is to prepare the appropriate solutions in the beaker and the test tube, then place both the beaker and the test tube in a water bath (larger beaker) containing either ice-water or warm water, depending on the desired temperature. Put the thermometer in the test tube. Allow the solutions to remain in the water bath until the correct temperature has been reached, then combine the mixtures and time the reaction as before. Make sure to record the temperatures in your lab notebook. CLEAN-UP: Dispose of the waste in the waste container in the fume hood. Clean up your glassware and return your equipment to the proper location. The 10 mL beakers, timer, stir bar, digital thermometer and reagents must be placed in the plastic container and returned to the equipment cart. Other glassware must be placed in your equipment drawers. Make sure to return the equipment to the same drawer from which you took it!

How to calculate the rate of reaction:

Using the dilution formula, calculate the concentration of S2O32- in each mixture. According to the stoichiometry of the clock reaction in Equation 4, the number of moles of I2 is one-half the the number of moles of S2O32-. What will be the concentration of I2 ([I2]) when the blue color appears? You will use this number as ∆[I2] in all of your rate calculations. Show this calculation in your lab report!!!

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How to calculate the Ion Molarities:  To determine the concentration of an ion in solution, you must consider the stoichiometric relationship between the ionic compound and the number of ions formed in solution. For example, a 0.20 M solution of KI releases 0.20 M I- and 0.20 M K+ ions. The situation would be different if the source of I- was CaI2, since 2 moles of I- ions would be released for each mole of CaI2 that dissociates.  Since you are mixing solutions, you need to take into account the dilution that occurs when the solutions are combined. That is, you need to calculate the concentration of each reactant after combining the solutions, but before the chemical reaction begins. Note that the total volume of each solution is 48 drops. Thus, in the dilution formula, M1V1 = M2V2, V2 is always 48 drops, and V1 is the number of drops of the individual solution added to the mixture.

Calculations and Results Instructions Before week 2, try to complete the following calculations and tables: 

 



  



Convert time for each trial to seconds, and calculate the average time in seconds for the first five runs. Also, calculate the rate of reaction (∆[I2]/∆t) for Runs 1-9 as explained above. Use the dilution formula to calculate the concentration of ions I- and S2O82- for all the runs as explained above. Use Runs 1, 2 and 3 to determine x, the order for I-. Pick two of these runs and use the initial rates method (shown in the example on page 2) to find the value for x. Pick a different combination among Runs 1, 2 and 3 to verify the value obtained for x. Calculate and determine the order for I-, rounding the value for x to the nearest whole number. Use Runs 2, 4 and 5 to determine y, the order for S2O82-. Pick two of these runs and use the initial rates method (shown in the example on page 2) to find the value for y. Pick a different combination among Runs 2, 4 and 5 to verify the value obtained for y. Calculate and determine the order for S2O82-, rounding the value for y to the nearest whole number. Substitute x, y, the rate of reaction, and your calculated ion concentrations back into the rate law (Equation 3) to calculate the rate constant, k, for runs 1-5 and 7-9. Run 6 is invalid because of Cu(NO3)2. Calculate the average rate constant, k, for the room temperature runs 1-5. Prepare “Results Table I” with rows for all nine runs and with columns for average time of reaction, the rate of reaction, concentration of I- in the mixture, concentration of S2O82- in the mixture, the temperature and k for runs 1- 5, 7, 8, and 9. This table must be completed for next week except the k column! Note: if you put the units in the column headings then you will not have to write the units in each cell in the table. Report the rate law in your results section by substituting in the values for the 3 constants: x, y, and k. Make sure to include units for your average room temperature rate constant.

During week two you will do the following:  Prepare “Results Table II” with rows for runs 3, 7, 8, and 9 and columns for: rate constant, ln k, 



temperature (in Kelvin), and 1 / T (what will the units be?). Use Graphical Analysis or Excel to prepare an Arrhenius plot with ln k on the y-axis vs. 1/T on the x-axis. (Note that "T" needs to be expressed in Kelvin rather than °C.) This plot should have four points, the original room temperature data for Run 3, and the runs you carried out at the three additional temperatures. Directions for GA: Make sure to enter the data in order from lowest temperature to highest temperature. Double-click on the graph and remove the check mark next to “connect points”. Select the regression line tool ("R=") from the toolbar at the top. From the “File” menu select “Page Setup”. Select “Landscape” orientation and click “OK”. To print a copy of the graph and data table, click on the printer symbol on the toolbar at the top. The “Printing Options” box will appear. Type your name and the names of your partners in the “Name” box , then click “OK”. If you use Excel make sure you print the graph with the straight line equation showing. Excel directions: http://web.gccaz.edu/~ksmith8/graph.htm

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

Use the Arrhenius plot you graphed to calculate the activation energy, Ea, in units of kJ/mol for this reaction. In an "Arrhenius plot", m = Ea/R, where the gas constant, R, is 8.314

J . K  mol

Questions for lab report: 1. For the three-step mechanism below, determine the following: a. Net chemical equation b. Rate law for each elementary step c. Overall rate law d. Molecularity of each step e. Identify the intermediate(s) and catalyst(s), if any. Step 1 – slow: ClO- + H2O  HOCl + OHStep 2 – fast: Br - + HOCl  HOBr + ClStep 3 – fast: OH- + HOBr  H2O + BrO2. The reaction of acetone with bromine is: C3H6O + Br2  C3H6OBr + HBr. The following data for initial rates was obtained. Determine the rate law, overall order for the reaction, and the rate constant (with units) for this reaction from the table below. Trial 1 2 3

[C3H6O] 0.10 M 0.20 M 0.10 M

[Br2] 0.10 M 0.10 M 0.30 M

Initial Rate (M/s) 1.64 x 10-5 1.65 x 10-5 4.93 x 10-5

3. Compare your average reaction time for Run 3 without Cu(NO3)2 to your reaction time with Cu(NO3)2. How did the addition of Cu(NO3)2 affect the rate of reaction? WHY did Cu(NO3)2 have this affect? What did Cu(NO3)2 do that changed the reaction time? 4. Compare your average reaction time for Run 3 at room temperature to your reaction times at the colder and hotter temperatures. What effect does changing the temperature have on the rate of reaction? (What happens at cold temperatures? What happens at hot temperatures?) Explain WHY temperature has this affect for both cold and hot temperatures?

Formal Lab Report – Use third person, past tense for reports! 1. 2. 3. 4.

5.

6. 7. 8.

Title, date, name, partners – can be a heading, doesn’t have to be a separate page. Purpose in complete sentences (not copied). Data measured in lab. Calculations: label each calculation, show the equation used, show the values you use in your sample calculation, and show the final answer as reported in your results. a. Show how you calculated the following (with units) for run 1: rate of reaction (show dilution and rate calculation), diluted concentrations of I- and S2O82-, orders of I- and S2O82-, k , and the average k for runs 1 – 5. Results: Include Results Table I and II, the rate law you determined, the average k for the room temperature runs, the Arrhenius plot showing the equation of the line, and the value of Ea in kJ/mol. Write your final rate law (i.e., including the average k (value and units) and order of each reactant) – there should be values for x, y, and k in your rate law. Questions 1 – 4; type the question and answer in complete sentences. Conclusion: Summarize and discuss your results from this experiment in several sentences (good length paragraph(s)). Also discuss at least 3 sources of error in your experiment. Restate the purposes and what results were obtained. Prelab initial page from each lab partner.

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