Experiment 2: The Rate of an Iodine Clock Reaction

Experiment 2: The Rate of an Iodine Clock Reaction Introduction: Some reactions, including most of the ones that you have seen before, occur so rapidly that they are over as soon as the reactants are mixed. There are many more reactions, however, that are slower; they may require minutes, hours, days, or even years to reach completion. A few chemical reactions are so slow that it is difficult to show that they happen at all. Why the wide differences in reaction rates? Some of the factors that affect the rate of a given reaction are the Activation Energy of the reaction, the concentrations of the reactants, the presence of a catalyst (which affects the Activation Energy), and the temperature at which the reaction occurs. Purpose: In the quantitative part of this experiment, you will determine the rate law for an iodine clock reaction and the influence of temperature on that reaction. There are actually many different iodine clock reactions. They all have a common feature: the completion of the reaction is signaled by the appearance of the dark blue-black color that is characteristic of the interaction of molecular iodine (I 2 ) with starch. This color may appear so abruptly that it is as startling as an alarm clock. The rate of an iodine clock reaction depends on the concentrations of the reactants, as you would expect. As a result, the time required for the appearance of the dark color can be adjusted by adjusting the concentrations of the reactants. In this reaction, the rate of the reduction of ammonium persulfate ((NH 4 ) 2 S 2 O 8 ) with potassium iodide (KI) will be studied. The net ionic equation for this reaction is: Rxn 1) S 2 O 8 + 2I - 2SO 4 + I 2 One of the purposes of this experiment is to determine the rate law for this reaction. The rate law's general form will be: Eq 1) Rate = k[s 2 O 8 ] q [I - ] r and you will need to determine values for the rate constant k and the exponents q and r. Remember that the exponents must be determined experimentally; they cannot be obtained simply by looking at the balanced equation. The rate law will be deduced by the initial-rate method, using three experiments. Experiment 1 will be the reference with which the other two will be compared. In Experiment 2, the initial concentration of the S 2 O 8 ions will be twice that in Experiment 1, but the concentration of I - ions will be unchanged. The general form of the rate law for each of these experiments will be: Experiment 1: Rate 1 = k[s 2 O 8 ] q [I - ] r Experiment 2: Rate 2 = k(2[s 2 O 8 ]) q [I - ] r = 2 q rate 1 Chem 401 Lab Manual M Samples Revision 1/30/15 page 3

The rate in the second experiment will be 2 q times the original rate. In Experiment 3, the initial concentration of S 2 O 8 ions will be the same as that in Experiment 1, but the concentration of I - ions will be doubled. The rate law will be: Experiment 3: Rate 3 = k[s 2 O 8 ] q (2[I - ]) r = 2 r rate 1 The rate in this experiment will be 2 r times the original rate. You can deduce the true values of q and r from the ratios rate 2 /rate 1 and rate 3 /rate 1 and the initial concentrations of S 2 O 8 and I - ions. After you have evaluated the exponents, you will know the general form of the rate law, but the value of the rate constant will be unknown. You can obtain it from the rate law and any one of the three rates. Each rate should yield the same rate constant within experimental error. However, in order to derive the exponents and the rate law, you need to measure the rate of the reactions. How? In an initial rates method, you need to know the initial concentrations of the reactants (just a dilution calculation), but you also must determine the concentration of at least one of the reactants during the reaction. In this experiment, we will indirectly measure the concentration of the persulfate ions (S 2 O 8 ). Based on the persulfate ion, the rate of the reaction is given by: Eq 2) Rate = S O 2 8 t To find the rate, you will need to know or measure Δ[S 2 O 8 ], the initial change in the concentration of S 2 O 8 ions, and t, the time elapsed during that change. An easy way to obtain Δ[S 2 O 8 ] is by coupling another reaction to the one we are studying. The new reaction is the reduction of I 2 by Na 2 S 2 O 3 (sodium thiosulfate). The net ionic equation for this reaction is: Rxn 2) I 2 + 2S 2 O 3 2I - + S 4 O 6 where S 4 O 6 is the tetrathionate ion. It is important to remember that this reaction is used only as a means of studying the rate of the first reaction. The new reaction is fast. As a result, I 2 is consumed in this reaction as fast as it is formed in the first reaction. Essentially there can be no I 2 present as long as there are S 2 O 3 ions present. However, these ions are being consumed because of the second reaction. As soon as all the S 2 O 3 ions have reacted, I 2 from the first reaction begins to accumulate in the solution. You can detect the sudden presence of I 2 by the sudden appearance of the dark color from the interaction of I 2 with starch, which will be used as an indicator. You will know Δ[S 2 O 8 ] at the instant that this dark color appears. How? The stoichiometry and relative rates of the coupled reactions require the following relationship between Δ[S 2 O 8 ] and Δ[S 2 O 3 ]: Eq 3) [S 2 O 8 ] = 0.5 [S 2 O 3 ] where Eq 4) [S 2 O 3 ] = final [S 2 O 3 ] - initial [S 2 O 3 ] Chem 401 Lab Manual M Samples Revision 1/30/15 page 4

Essentially all the S 2 O 3 ions will have reacted when the dark color appears, so the final concentration of these ions is essentially zero, and Δ[S 2 O 3 ] becomes: Eq 5) [S 2 O 3 ] = - initial [S 2 O 3 ] = -[S 2 O 3 ] 0 Therefore, Δ[S 2 O 8 ] can be found as: Eq 6) [S 2 O 8 ] = -0.5 [S 2 O 3 ] 0 By combining Equation 2 and Equation 6, we obtain: Eq 7) rate rxn = 0.5 S O 2 3 Δt 0 If we keep the initial concentration of S 2 O 3 ions very small, Δ[S 2 O 3 ] will be small and Δ[S 2 O 8 ] will be even smaller. As a consequence, there will be little change in the concentration of the reactants during the elapsed time Δt. This is a necessary condition for the initial rates method. The time required for the appearance of the dark color is affected by the concentration of the reactants, as you would expect from the general form of the rate law. The time is also affected to a degree by the overall concentration of ions and their charges. This influence, which is well known, cannot be predicted from the rate law. You will maintain a constant volume by the addition of DI water as required. Although this is not quite correct, as we should also be maintaining a constant concentration of charges and ions, this is an advanced topic that we shall ignore now. Lastly, you will demonstrate the influence of the temperature on the rate of this reaction. To do this, you will redo Experiment 1 at three higher temperatures. The rate law will have the same general form at the higher and lower temperatures. As a result, q and r will not change when the temperature is varied. A new rate constant, however, will need to be found at the higher temperatures. The activation energy for the reaction can then be calculated via the Arrhenius equation using all four points and thru the point Adjusted form of the Arrhenius equation using two points. Chemicals: 0.20 M KI 0.20 M (NH 4 ) 2 S 2 O 8 (ammonium persulfate) 3% Starch Sln distilled water 0.010 M Na 2 S 2 O 3 (sodium thiosulfate) Day 1 Equipment: Thermometer 16x150 test tubes, 3 & test tube rack Waste Beaker 2 ml Pipetman Pipet Tips, 5 Stopwatch 10 ml Beakers, 5 Stir rods, if desired Day 2 Equipment: Thermometer 20x150 mm Test Tubes, 5 & large test tube rack Waste Beaker 2 ml Pipetman Pipet Tips, 5 Stopwatch 10 ml Beakers, 5 Hot Water Baths Stir rods, if desired Chem 401 Lab Manual M Samples Revision 1/30/15 page 5

Waste Disposal: All solutions must be disposed of in the proper waste container. Day 1 Procedure: 1. Label 3 clean, dry test tubes as follows: Rxn 1, Rxn 2, Rxn 3. These are the reactions you will conduct on Day 1. Obtain a tray to work on. 2. Label 4 10 ml beakers with the different chemical formulas; DI Water; KI; Starch; Na 2 S 2 O 3. Make sure they are clean & dry. Try to pour around 3-4 ml of each chemical into the appropriately labeled beaker. BE CAREFUL TO POUR THE CORRECT CHEMICAL INTO YOUR BEAKERS! READ AND REREAD THE LABELS! 3. Note and record the temperature of the lab to the nearest 0.1 C. 4. Make sure all your pipet tips are clean & dry. One of these tips is dedicated to ammonium persulfate. If you have 5 pipet tips, label them carefully with the 5 different chemicals, and place each tip in the beaker with the appropriate chemical (the 5 th tip is for ammonium persulfate, Step 8 below). 5. Using a pipetman, add the appropriate amounts of KI to the labeled reaction test tube. BE CAREFUL TO USE THE CORRECT CHEMICAL AND CORRECT AMOUNT! Constantly check your Table of Volumes. 6. Clean & dry the pipet tip if you have to reuse it for a different chemical. Otherwise, place it back in the KI beaker. 7. Repeat Steps 5 and 6 for DI Water, Na 2 S 2 O 3, and 3% Starch. Mix the contents of the reaction test tube by agitating by hand or by using a stir rod. If you use a stir rod, clean and dry it each time after using it. 8. Pour about 3-4 ml of ammonium persulfate, (NH 4 ) 2 S 2 O 8, into a clean, dry, labeled 10 ml beaker. You will start each reaction upon the addition of this chemical. 9. READ STEPS 10 AND 11 COMPLETELY BEFORE PROCEEDING! 10. Start the stopwatch JUST as you pipet the required amount of (NH 4 ) 2 S 2 O 8 to reaction test tube 1. Mix thoroughly, and you may continue to shake it gently or stir it as the reaction progresses. 11. Stop the stopwatch at the EXACT time when the solution turns dark (blue) throughout the solution. BE ALERT! This color may appear suddenly throughout the entire solution. (Note: Experiment 1 should take around 5-12 minutes depending on the exact solution concentrations.) 12. Record the elapsed time. 13. Repeat Steps 10-12 for Experiments 2 and 3 in test tubes 2 and 3, using the appropriate amounts of (NH 4 ) 2 S 2 O 8. Check your data with your instructor to see if you need to repeat any Experiment. 14. Dispose of all waste in the appropriate container, and clean all glassware and the tray. Return any borrowed items (Except the pipetman, keep reading!). Clean your area. Take your pipetman along with your data to the instructor and obtain a data initial. Put the pipetman back on the pipetman rack. 15. Do Data Analysis Steps 1 and 2 and obtain a data analysis initial. Chem 401 Lab Manual M Samples Revision 1/30/15 page 6

Day 2 Procedure: On day 2, you will need approximately 2 ml KI, DI Water, Starch, and Na 2 S 2 O 3 ; and 1 ml of (NH 4 ) 2 S 2 O 8. Label 10 ml beakers and add the appropriate amount of each chemical (less than half full). Make sure these beakers are clean! You may do Experiments 4-7 in any order, although the procedure is written for a specific order. 1. Label 4 20x150 mm clean and dry test tubes as follows: 4, 5, 6, 7. These are the 4 reactions you will conduct on Day 2. The 5 th test tube is dedicated for ammonium persulfate. Label the 5 th test tube (NH 4 ) 2 S 2 O 8. Obtain a spill tray. 2. Make sure the pipet tips are clean and dry or you will need to rinse them and prime them with the chemical. As on Day 1, add the appropriate amount of KI, DI Water, Na 2 S 2 O 3, and 3% Starch to each of your 4 labeled reaction test tubes. Mix well by shaking by hand or by stirring with a stir rod. If you use a stir rod, remove the stir rod and clean and dry it each time you use it. 3. For Experiment 4, have the stopwatch ready. Pipet 0.300 ml of ammonium persulfate into the ammonium persulfate test tube and then pour it into your waste beaker (this way we coat the bottom of the test tube). Again, pipet 0.300 ml of ammonium persulfate into the ammonium persulfate test tube and go to step 4. 4. Pour the contents of the ammonium persulfate test tube into test tube #4 (Experiment 4), start the stopwatch, mix quickly by shaking by hand, and place the reaction test tube back in the test tube rack.* Note: Put the ammonium persulfate test back in your test tube rack at your bench. You may reuse the ammonium persulfate test tube for Experiments 5-7 without cleaning and drying it. 5. When the reaction is complete (turns blue as on Day 1), stop the stopwatch, and record the elapsed time. 6. For Experiment 5, you will use a hot water bath at approximately 40 C. 7. Place this labeled test tube in the hot water bath. 8. Pipet 0.300 ml of the (NH 4 ) 2 S 2 O 8 into the ammonium persulfate test tube. Place this labeled test tube in the hot water bath as well. 9. Let both test tubes sit in the hot water bath for 5 minutes. 10. GET READY TO START TIMING! READ THE STEPS 11-13 BEFORE PROCEEDING! 11. Remove both test tubes from the hot water bath (to avoid contaminating the hot water bath), quickly add the warm (NH 4 ) 2 S 2 O 8 to the warm reaction test tube. Start the stopwatch. 12. Shake the reaction test tube for a few seconds (without spilling or splashing). Replace the reaction test tube in the hot water bath. The ammonium persulfate test tube may go back in your test tube rack at your bench. 13. BE ALERT! Stop the stopwatch at the EXACT time when the solution turns dark (blue). 14. Remove the test tube, and note and record the actual temperature of the hot water bath. 15. Record the elapsed time. 16. Repeat steps 6-15 for Experiment 6 at approximately 50 C and for Experiment 7 at approximately 60 C. 17. Dispose of all waste in the appropriate container, and clean all glassware and the tray. Return any borrowed items (Except the pipetman, keep reading!). Clean your area. Take your pipetman along with your data to the instructor and obtain a data initial. Put the pipetman back on the rack. 18. Continue working on the Data Analysis and obtain a data analysis initial at the end of the lab. * This seems awkward, but you are mimicking what you will be doing with the other reactions this day. Chem 401 Lab Manual M Samples Revision 1/30/15 page 7

Data: Fill in the data section at the end of the lab. Data Analysis: 1. Calculate [S 2 O 3 ] 0 for the reaction solution for experiments 1-3. Remember that these are diluted solutions, so you are solving for M 2 in the dilution equation M 1 V 1 =M 2 V 2. M 2 is the initial concentration just before the reaction started but after all of the solutions were mixed together. Look at the volume table on page 9 so that you use the correct start concentrations, the correct volume of S 2 O 3, and the correct total volume for all of the experiments. Show all units and work and equations. 2. Calculate the initial rates for Exp. 1, 2, and 3 from Equation 7 and the [S 2 O 3 ] 0 you just calculated in Data Analysis 1. 3. Calculate [S 2 O 8 ] 0 and [I - ] 0 for each experiment (including Exp 4-7) using M 1 V 1 =M 2 V 2. Show all units and work. Again, look at the volume table on page 9 so that you know exactly what volumes, etc. were actually used for each experiment. 4. Make a Table like the following (fill in the blanks with your results for Exp 1-3): Exp # [I - ] 0 [S 2 O 8 ] 0 rate (M/s) 1 2 3 Determine the reaction orders with respect to I - and S 2 O 8. These should be whole numbers, so you will need to round to the nearest whole number. Show all units and work, however, you may be able to determine the reaction orders by eyeballing the start concentrations and the rates. 5. Write the correct rate equation. You are now done with Day 1 data! In Steps 6 and 7 you will use Day 2 data and this rate law to calculate the rate constant k for 4 different temperatures. 6. Calculate the rate for Experiments 4-7 exactly as you did for Experiments 1-3 in Steps 1 and 2 above. 7. Calculate the rate constant k for Experiments 4-7 using the rates you just calculated in Step 6, the concentrations of [S 2 O 8 ] 0 and [I - ] 0 from Step 3, and the rate equation you found in Step 5. Show all units and work. 8. Using your room temperature k (from Exp 4 on Day 2), and one of your k values from an elevated temperature, use the adjusted point form of the Arrhenius Equation to calculate E a. Show all units, work, and equations. 9. Now you have 4 rate constants for 4 different temperatures. Determine the Activation Energy from the slope of a linear plot of lnk vs. 1/T. Y-axis is ln(k). Show all units, work, and equations. Discussion/Errors: 1) Compare your rate law to the known room temperature rate law of rate = 0.0040[I - ][ S 2 O 8 ]; 2) discuss how your k values change with temperature; 3) calculate the %- error between your room temperature k and the accepted room temperature k of 0.0040; 4) compare your two E a values (which should be more accurate and why); and 5) discuss possible or known sources or errors. Use your graphs to highlight errors that may be seen; graphs often show errors in accuracy and precision. Chem 401 Lab Manual M Samples Revision 1/30/15 page 8

Data/Observations: Pipetman Day 1: Pipetman Day 2: Laboratory Temperature: Day 1 3% Volumes (ml) of Solutions for Experiments (May Change): Experiment 0.20 M DI Water 0.010 M 0.20 M KI Na 2 S 2 O 3 Starch (NH 4 ) 2 S 2 O 8 1 0.500 0.500 0.500 0.500 0.300 2 0.500 0.200 0.500 0.500 0.600 3 1.000 0 0.500 0.500 0.300 4, 5, 6, 7 0.500 0.500 0.500 0.500 0.300 Note that the total volume for each experiment is 2.300 ml (May change). Elapsed Times for Iodine Clock Reactions Experiment Elapsed Time (s) Observations (what color blue, etc.) 1 2 3 4 (Day 2) 5 (Day 2) 6 (Day 2) 7 (Day 2) Temperature of Reaction Solution for Experiment 4: 5: 6: 7: Formulas: Arrhenius Equation (in form of equation of line: y = mx + b) ln(k) = E a RT + ln(a) Pt Adjusted Form of Arrhenius Equation: ln k 2 k 1 = E a 1 1 R T 1 T 2 Chem 401 Lab Manual M Samples Revision 1/30/15 page 9