Chemistry 134 � Lab 7 The Kinetic Study of the Iodination of by akimbo

VIEWS: 1,701 PAGES: 6

									Chemistry 134 – Lab #7 The Kinetic Study of the Iodination of Acetone Goals • To determine the order of the reaction with respect to acetone, iodine, and hydrochloric acid. • To determine k at an assigned temperature. • To determine the activation energy of the reaction, using data from classmates. Background The Differential Rate Law. For the general reaction (1), the rate is expressed as the change of concentration with respect to time, t, of any reactant or product. (1) aA + bB dD + fF  1  ∆[A]  1  ∆[F]  1  d [A]  1  d [F] Rate = −  =  or Rate = −  =  (2)  a  ∆t  a  dt  f  ∆t  f  dt Since A is consumed and F is produced by (1), their rates (derivatives) have opposite signs. Thus we introduce a minus sign in front of ∆[A]/∆t to ensure a positive rate. Also, unless a=f, -∆[A]/∆t and ∆[F]/∆t are different. In general, the initial rate of reaction is a function of the concentrations of reactants, as indicated in (3).  1  ∆[A]  1  ∆[F] Rate = −  =  = k [ A] m [ B] n (3)  a  ∆t  f  ∆t k is the rate constant of the reaction, and [A] and [B] are the molar concentrations of reactants and products in (1). The exponents m and n appearing in (3) are usually either positive integers or zero. Because (3) is a differential equation, it is called the differential rate law for (1). The goal of a kinetics study is to determine the differential rate law for the reaction (i.e., the values of the exponents m and n) and the numerical value of the rate constant k at each of several temperatures. Suppose that appropriate kinetics experiments carried out on (1) give m=2 and n=1. The differential rate law for the reaction is: (4) Rate = k [ A]2 [ B] The exponents m and n are called orders. We say that the reaction is second order in A and first order in B. It is important to realize that there is no necessary connection between the reaction orders m and n in (3) and the stoichiometric coefficients a and b in (l). The Method of Initial Rates. A straightforward way to determine experimentally the specific form of the rate law (3) is to measure the initial rate, -(∆[A]/∆t)0 , as a function of each of the initial concentrations [A]0 , [B]0 , etc, and from the data deduce the reaction orders. Suppose that we found that -(∆[A]/∆t)0 quadrupled when we doubled the [A]0 while keeping the other concentrations constant. We would conclude that m=2. Similarly, if -(∆[A]/∆t)0 decreased by a factor of 2 when [B]0 was cut in half while keeping other concentrations the same, we would conclude that n=1.

It may be difficult to deduce the reaction order when the value is not integral. The reaction order may be calculated by forming a ratio of the two trials: n (5) rate1 k[ A]m [ B]1 1 = rate2 k[ A]m [ B] n 2 2 If the concentration of B is the same in the two trials, then (5) simplifies to:
m rate1 [ A]1 rate1  [ A]1  = =  m or rate2 [ A] 2 rate2  [ A] 2  Finally, we take the log of both sides and then solve for m:  rate1  log    rate1   [ A]1   rate2  log   = m log   ; m=  rate2   [ A] 2   [A]1  log   [ A] 2  m



Calculation of Activation Energies. The relationship between the reaction constant, k, and the temperature is given by the equation: Ea (8) − E 1 k = Ae RT or ln k = − a + ln A R T where Ea is the activation energy, R is the gas constant, T is the temperature, k is the rate constant, and A is a constant that depends on the steric requirements of the reaction. The second form of (8) is a linear equation. The activation energy may be determined by plotting ln k vs. 1/T, which yields a line of slope –Ea/R. Recall that the slope of a line is the change in the y coordinate divided by the change in the x. Substituting in our m, x, and y variables yields: E ln k 1 − ln k 2 − a = 1 1 (9) R T1 − T2

The Experiment. Today we will examine the reaction between acetone (CH3 COCH3 ) and iodine, in the presence of an acid catalyst (HCl). The overall reaction may be written as:


This overall reaction does not tell us about how the reaction actually proceeds, only what the final products will be. To study this reaction, we will need a way to measure the rate of change of the concentration of a reactant or product. Because I2 (aq) is yellow, it provides a convenient way to monitor the reaction. The rate of this reaction will be ∆[ I 2 ] Rate = − (11) ∆t You will have a chance to verify in this experiment that the rate with respect to the iodine is zero. During this experiment, you will use relatively high concentrations of acetone and

hydrochloric acid, and low concentrations of iodine. The result is that the reaction is pseudozero order, because the amounts of the acetone and acid are relatively constant during each trial. Thus, during the time while the iodine is being consumed, the rate will be constant. You will measure the time that it takes for all of the iodine to be consumed by the reaction, and use that as ∆t in the rate expression. The change in [I2 ] is given by: ∆[I2 ] = 0 - [I2 ]o Materials 1000 mL beaker 3 100 (or 250) mL beakers Two similar test tubes 250 mL Erlenmeyer flask Ring stand and clamps Thermometer Hot plate/magnetic stirrer Magnetic stir bar Ice (if assigned a temperature lower than 25°C) Stopwatch 3 graduated pipets (25 mL) 4.0 M acetone (aq) 1.0 M hydrochloric acid (aq) 0.005 M I2 in 0.05M KI (aq) Procedure Setup. Fill a 1000 mL beaker with water, and adjust the temperature of the water to your assigned temperature by heating on the hot plate or adding ice, as appropriate. Use the magnetic stir bar to keep the temperature uniform throughout the beaker. During this experiment, you will need to monitor the temperature, and make adjustments with heat or ice as necessary to maintain the assigned temperature. Obtain two identical test tubes. Clean and rinse thoroughly, and then fill both tubes with distilled water. Hold the test tubes up to a piece of white paper and compare the color. Both should appear the same. (If they do not, try a different pair of test tubes.) Keep one tube filled with water; this is your reference test tube. Empty the other tube, which you will use for the reactions. Pour approximately 60 mL of the following solutions into clean, labeled beakers: 4.0 M acetone, 1.0 M HCl, and 0.005M I2 . Cover each beaker with a watch glass to reduce evaporation.
5 6 4 7 5 6 4 7 3 2 8 3 2 8 9 9 1 1 1 1 1 0

First set of concentrations. For your first trial, you can use 10.0 mL of each of the three reactants, and 20.0 mL of water. (Students working below 20°C: Use 20.0 mL HCl, 10.0 mL acetone, 10 mL I2 , and 10.0 mL water instead.) Pipet the water, the HCl, and the acetone in a 250 mL Erlenmeyer flask. Place the iodine in one of your test tubes. Put the Erlenmeyer flask and the test tube with the iodine into the water bath, as shown in the diagram. (The test tube can float in the bath, as long as it is big enough not to tip over.) Wait 5 minutes for the glassware

and contents to reach the temperature of the water in the bath, which should still be the assigned temperature. After the five minute temperature equilibration, remove the test tube and Erlenmeyer from the water bath. Begin timing the reaction as you pour the contents of the test tube into the Erlenmeyer. Swirl the Erlenmeyer briskly to mix the contents, then pour part of the mixture back into the test tube. Put the test tube back into the water bath. Watch closely. When the yellow color of the iodine is almost gone, pull the test tube from the water bath and hold it up again the white paper, next to your reference test tube. When the color of the reaction test tube matches the color of the reference test tube, stop timing. Record the time (in seconds) in your data table. Rinse the Erlenmeyer and the reaction test tube thoroughly. Now do a second trial, using the same amounts of each reagent. Compare the two times. If they differ by more than 10%, run a third trial. Calculate the average time for the reaction, and then calculate the rate of the reaction using equation (11). Changing the concentrations. After you have determined the rate for one set of concentrations, it is time to study how the rate varies with changes in the concentrations. What changes should you make to allow you to learn how the acetone concentration affects the rate? Plan a new reaction mixture, which allows you to study just the effects of changing the acetone concentration. You should keep the total volume constant, so you will need to adjust the amount of water you use in the reaction mixture. If you’re not sure if your conditions will allow you to learn about the order of acetone, check with the instructor before proceeding. Perform at least two trials, and record the times. Calculate the rate for this new set of conditions, and then calculate the order of the reaction with respect to acetone. Next, plan some reaction conditions that will allow you to determine the order with respect to the hydrochloric acid, and calculate the order with respect to hydrochloric acid. Again, perform two or more trials, and calculate an average ∆t. Finally, plan and perform reactions to determine the order with respect to I2 . We have been assuming that the rate with respect to I2 is zero, but we should verify this. (NOTE: When calculating the rate of this reaction, don’t forget that the concentration of iodine has changed!) Procedural hints • If you were assigned a low temperature, you may want to choose sets of conditions that will cause the times to be shorter. Increasing the concentration of acetone or hydrochloric acid will decrease the reaction time. However, because we determine the time required for all of the iodine to be consumed, adding more iodine means that the reaction will take longer. If your reaction times are long, halve the amount of iodine instead of doubling it. • If you were assigned a high temperature, your reactions will be FAST! Doubling the concentrations of acetone or the acid will make the reactions even faster, and may make it difficult to accurately determine the reaction times. Halve the concentrations instead. • Regardless of temperature, make sure you adjust the amount of water used so that your total volume is always 50.0 mL. Calculations. If you haven’t already done so, calculate the order of the reaction with respect to I2 , acetone, and HCl. If your calculated orders are close to being whole numbers, round to the nearest integer. Use the rounded orders for the rest of the calculations. Write the rate law using

the rounded orders. If the orders you calculate are very far from whole numbers, consult the instructor. Calculate the average k value for your four reactions. Report your average k value and assigned temperature to the rest of the class, and record the class k values and temperatures in a data table. Finally, you will determine the activation energy (Ea) of this reaction. To do so, calculate ln(k) and 1/T for each data point from the class. Record these values in your data table. Make a graph of ln k vs. 1/T. (ln k belongs on the y axis.) Draw the best straight line through the data. Use the graph to calculate Ea, as described in the introduction. Pre-lab preparation. A detailed description of the basic procedure should be written in your lab notebook. Make a data table that lists the volume of each reactant (and water) that you will use for your trials. Leave room to record up to three times for each set of conditions. Make a second data table which shows the molarity of each reactant for each trial, and the rate for each trial. Finally, make a data table for the lnk vs. 1/T plot. Sample data tables acetone (mL) 10 HCl (mL) 10 I2 (mL) 10 H2O (mL) 20 Times (s) (∆t) average time


(choose the volumes that go here)

[acetone] (M) I II III IV

[HCl] (M)

[I2] (M)

rate ([I2]/average ∆t) (M/s)

T (°C) T (K) 1/T (1/K) k (Mx/s)* ln k 10 … … … … 40 *The units for k will depend on the order of the reaction. Put the correct units in your data table after you determine the order of the reaction.

Post-lab 1. Please summarize your results. Report your rate law with the experimentally-determined orders, your value for k, and your value for the activation energy. 2. No one in the class was assigned a temperature below 10°C or above 40°C. Using your data on the relationship between k and T, predict the value of k and ∆t for the reaction at 0°C and 50°C, under the conditions listed below. (Show work.) T (°C) 0 50 acetone
mL M

mL M

I2 (mL)
mL M

mL M

expected k

expected time (∆t) (s)

10 10

10 10

10 10

20 20

3. If you wanted to study k below 10°C or above 40°C, what changes would you need to make to the experiment protocol?

Catherine Sarisky, 3/01. Figures created with ACD/Chemsketch.

To top