Chem 1B Dr. White 175 Experiment 26 - Kinetics Objectives To determine the rate law for the reaction between iodide and bromate under acidic conditions To investigate the effect of temperature on rate To determine the activation energy for the the reaction between iodide and bromate under acidic conditions Introduction This experiment involves the study of the rate properties, or chemical kinetics, of the following reaction between iodide ion and bromate ion under acidic conditions: 6 I - (aq) + BrO 3 - (aq) + 6 H + (aq) 3 I 2(aq) + Br - (aq) + 3 H 2O(l) (Rxn 1) This reaction proceeds at an easily measurable rate that depends on the concentrations of the I -, BrO 3 -, and H + ions according to the rate law. For this reaction, the rate law takes the form rate = k [I - ] x [BrO 3 - ] y [H + ] z (1) One of the main purposes of the experiment will be to evaluate the rate constant, k, and the reaction orders x, y, and z for this reaction. We will also investigate the manner in which the reaction rate depends on temperature and will evaluate the activation energy, E a, for the reaction. Our method for measuring the rate of the reaction involves what is frequently called a "clock" reaction. In addition to 1, whose kinetics we will study, the following reaction will also be made to occur simultaneously in the reaction flask: I 2(aq) + 2 S 2O 3 2- (aq) 2 I - (aq) + S 4O 6 2- (aq) (Rxn 2) As compared with 1, the rate of 2 is essentially instantaneous. The I 2 produced in 1 reacts completely and instantaneously with S 2O 3 2- so that until the S 2O 3 2- is used up the concentration of I 2 is effectively zero. As soon as the S 2O 3 2- is used up, the I 2 remains in solution and is made evident by the reaction with starch. The reaction of I 2 with starch produces a blue color. The appearance of the blue color tells us when a given amount of BrO 3 - has reacted. We will not calculate the actual amount of BrO 3 - that has reacted, but use an arbitrary value of 1000. This will give us the relative rate which is equal to 1000/t, where t is the time it takes to turn blue. concentration conditions. Measured amounts of each of these ions in water solution will be mixed in the presence of a constant small amount of S 2O 3 2-. The time it takes for each mixture to turn blue will be measured. The time obtained for each reaction will be inversely proportional to its rate. By changing the concentration of one reactant and keeping the other concentrations constant, we can investigate how the rate of the reaction varies with the concentration of a particular reactant. Once we know the order for each reactant, we can determine the rate constant for the reaction. This is illustrated in the example below. Example: The kinetics of the following reaction was investigated: 2NO (g) + O 2 (g) 2NO 2 (g). The following data were collected: Exp. [O 2] (M) [NO] (M) Relative Rate (M s - 1 ) 1 0.0010 0.0010 7.1 2 0.0040 0.0010 28.4 3 0.0040 0.0030 255.6 To find the rate law for the reaction we need to use two experiments where the concentration of one reactant is held constant. In experiments 1 and 2 the concentration of NO is constant and the concentration of O 2 is changed. Thus using experiments 1 and 2 we can find the reaction order with respect to O 2: rate 2 = k[o 2 ] x y 2 [NO] 2 rate 1 k[o 2 ] x y 1 [NO] = k(0.0040 M) x (0.0010 M) y 1 k(0.0010 M) x (0.0010 M) y 28.4 M s 1 0.0040 M = 1 7.1 M s 0.0010 M 4 = 4 x x = 1 x A similar process is used to determine the rate order with respect to NO: rate 3 = k[o 2 ] x y 3 [NO] 3 rate 2 k[o 2 ] x y 2 [NO] = k(0.0040 M) x (0.0030 M) y 2 k(0.0040 M) x (0.0010 M) y 255.6 M s 1 0.0030 M = 1 28.4 M s 0.0010 M 9 = 3 y y = 2 Now that we know x and y, we can use the rate law to determine the rate constant: y In this experiment, we will carry out the reaction between BrO 3 -, I -, and H + ions under different
176 Chem 1B Dr. White rate = k[o 2 ][NO] 2 Table 1: Volumes and reagents for Flask I rate (125 ml) k = [O 2 ][NO] 2 Volume of Volume of 0.0010 M H2O (ml) Using experiment 1, 7.1 M/s k = [0.0010 M][0.0010 M] 2 = 7.1x109 M -2 s -1 In the second part of the experiment we will investigate how the rate of the reaction depends on temperature. In general the rate increases sharply with temperature. By measuring how the rate varies with temperature we can determine the activation energy (E a), the minimum energy required for the reaction to occur, by making use of the Arrhenius equation: ln k = E a R 1 + ln A T (2) In this equation, k is the rate constant at the Kelvin temperature T, E a is the activation energy, A is another constant for a given reaction and R is the gas constant. By plotting ln(k) against 1/T we should obtain, by Equation 2, a straight line whose slope equals - E a/r. From the slope of that line, we can easily calculate the activation energy. Procedure Part 1: Dependence of Rate on Concentration See the tables below. Except for the indicated amounts of solutions used, the 5 trials will follow the same procedure. 1. Use the following procedure for each trial. Between trials, rinse the flask with DI water and shake out the extra water. a. Measure into the 2 different flasks (Flask I and Flask II) the reagents shown in the tables. Use graduated cylinders to measure all the reagents. b. Have a thermometer and a stopwatch ready. Pour the contents of Flask II into Flask I and immediately start the stopwatch. Swirl the reaction flask. The instant the mixture turns blue, stop the stopwatch. Record the elapsed time of the run in seconds. Measure and record the temperature. Volume of 0.010 M KI (ml) Na2S2O3 (ml) 1 5 5 5 2 10 5 0 3 5 5 0 4 5 5 0 5 4 5 6 Table 2: Volumes and reagents for Flask II (125 ml) Starch Indicator Volume of 0.040 M KBrO3 (ml) Volume of 0.10 M HCl (ml) 1 5 5 3 drops 2 5 5 3 drops 3 10 5 3 drops 4 5 10 3 drops 5 2.5 7.5 3 drops Part 2: Dependence of Rate on Temperature 1. In this part of the experiment, use reaction mixture 1 and follow the instructions above. However, the reaction mixtures will be at the following temperatures: about 0 C, about 10 C, and about 40 C. Record the actual temperature for each trial. a. 0 C: Use ice to cool your reaction flasks to 0 C. Leave the flasks in the ice for about five minutes before you start the trial and during the reaction. b. 10 C: Use a plastic tray with cool water to get the solutions to about 10 C. Leave the flasks in the water for about five minutes before you start the trial and during the reaction. c. 40 C: Use the one of the thermostatted baths set up for the trial at 40 C. Leave the flasks in the water for about five minutes before you start the trial and during the reaction. 2. Prepare an Arrhenius plot on Excel so that the activation energy can be determined.
Name: Chem 1B Dr. White 177 Lab Day/Time: Experiment 26: Kinetics Data and Results Part 1: Rate Law 1 2 3 4 5 Time, t, for Color Change (s) Relative Rate (1000/t) [I - ] (M) [BrO3 - ] (M) [H + ] (M) Temperature ( C) Show how you determined the [I - ] for 1 (it is diluted from the given stock solution by mixing all the reagents together): Show how you determined the [BrO3 - ] for 1 (it is diluted from the given stock solution by mixing all the reagents together): Show how you determined the [H + ] for 1 (it is diluted from the given stock solution by mixing all the reagents together): Determine the value of the reaction order x. Show your work below. Determine the value of the reaction order y. Show your work below.
178 Chem 1B Dr. White Determine the value of the reaction order z. Show your work below. Write the General Rate Law: Determine the Rate constant for each trial and fill in the following table (include the units of k): 1 2 3 4 5 Average Show your calculation for trial 1 below: k ( ) Briefly explain why k should have nearly the same value for each of the reactions. For 5, use the average k and the appropriate concentrations to predict (calculate) the relative rate. Then use this value to predict (calculate) the reaction time, t, for 5. Show your calculations below. relative ratepredicted tpredicted tobserved
Chem 1B Dr. White 179 Calculate the percent difference of your reaction times. % Difference = Part 2: Temperature Effects Approximate Actual Temperature Temperature ( C) ( C) Time t for Color Change (s) Your Data: 0 C 10 C 20 C* 40 C *Data from 1 in part 1. Relative Rate (1000/t) k ln k 1/T (K - 1 ) Attach a copy of your Arrhenius Plot made in Excel. Line of best fit from your Arrhenius Plot: Determine the activation energy (in kj/mol) and show your calculation below.
180 Chem 1B Dr. White Experiment 26: Kinetics Postlab Questions 1. A student studied the clock reaction described in this experiment. She set up a reaction mixture by mixing 10.0 ml of 0.010 M KI, 10.0 ml of 0.0010 M Na 2S 2O 3, 20.0 ml of 0.040 M KBrO 3 and 10.0 ml of 0.10 M HCl using the procedure given. It took 40.0 seconds for the color to turn blue. a. She found the concentration of each reactant in the reacting mixture by realizing that the number of moles of each reactant did not change when that reactant was mixed with the others, but that its concentration did. The volume of the mixture was 50.0 ml. Find the concentration of each reactant. [I - ] = M; [BrO 3 - ] = M; [H + ] = M b. What is the relative rate of the reaction (1000/t)? c. The student did a second trial in which she changed the amount of one reagent. She mixed 10.0 ml of 0.010 M KI, 10.0 ml of 0.0010 M Na 2S 2O 3, 10.0 ml of 0.040 M KBrO 3, 10.0 ml of 0.10 M HCl and 10.0 ml of H 2O. It took 81.0 seconds to turn blue. Find the concentration of each reactant after mixing: [I - ] = M; [BrO 3 - ] = M; [H + ] = M d. What is the relative rate of the reaction in the second trial(1000/t)? e. Use the information from the two trials above to determine the order with respect to BrO 3 -. n = (nearest integer)