Three Methods for Studying the Kinetics of the Halogenation of Acetone James P. Birk and David L. Walters Arizona State University, Tempe, AZ 85287-1604
The halorrenation of acetone has been used for a long time to introduce students to kinetic studies (1, 2). when iod~neis mixed with acetone, the following reaction occurs.
H O H
I II I H--C-C-C-H I I H
H
+ 12
-
H O I
I I
H-C-C-C-H H
II
I I
+ HI
H
An analogous reaction occurs between bromine and acetone. The reaction is quite slow in the pH range of 4 7 , but it occurs quite rapidly if the pH is below 3 or above 8 (3). This implies a n acid or base dependence. Assuming the iodination of acetone follows simple kinetics, the rate law has the following form. Figure 1. Absorbance of iadine in solution plotted as a function of iodine concentration for the Blocktronic colorimeter. Of all the participants in this reaction, iodine (or bmmine) is the only substance that gives a colored aqueous solution. The color intensity can be used to measure the amount of halogen left in solution. Measuring the absorbance of a reacting solution over time can be used to determine the rate of reaction. We wish to report on three methods for carrying out a kinetic study of this reaction. They differ primarily in the device used to determine color intensity in reaction mixtures: the SERAPHIM Blocktronic colorimeter (41, a commercial spectrophotometer, and the human eye (1,2). We will use various factors to compare the methods: the accnracy and precision, the cost effectiveness, and the ease with which the student can perform the kinetic study. Blocktronic Colorimeter First we describe the experiments carried out with a Blocktronic colorimeter connected to an Apple IIe computer, using SERAPHIM software (5)to calculate the absorbance. Since this colorimeter is more sensitive to the color of iodine than to the color ofbmmine, iodine was used in these experiments, in the form of potassium triodide solutions.' Since the Blocktronic does not measure monochromatic light, it was necessary to first determine that the absorbance is directly proportional to the concentration of iodine ([I2]+ [I3-]).Measurements were made relative to a water blank on four different solutions, which contained different concentrations of iodine in the range of 1.67 x lo3 M to 9.95x M. The solutions also contained 0.32 M HCI to give an ionic medium similar to that used in the reaction mixtures and to prevent disproportionation of 12.
The iodine stocksolution was made bvdissolvina 0.506a of I, and in 100 mL of distilled water, giving iconcei;trati& of
1.603 g of KI 0.02 M I* + 13-.
Experimental The reaction mixtures were made up with both the acetone and hydrogen ion concentrations in large excess to the iodine concentration, so they remained virtually constant during the reaction. Acetone solution, hydrochloric acid, and distilled water were first mixed together in the test tube, which was used as a cnvette. The iodine solution was then added to initiate the reaction. The conten- of the tcst tube were quickly mixed with a class stirrine rod. Then thev were olaced in the Blocktmnit colorimeter in which the absorbance of the solution was measured over time until the reaction was over. ARer the reaction reached completion, the temperature of the solution was measured. Aleasbsquares fit of the data to a straight line, shown in Figure 1, gave a correlation coefficient of R2 = 1.00 for the following relationship between absorbance and concentration.
-
-
Data Analysis Absorbance readings were converted to iodine concentrations. A plot of iodine concentration vs. time yields a straight line, as shown in Figure 2. The linearity of this plot indicates that the reaction is zero-order in [I21 ( b = 0). Least-squares analysis of the data, always with a correlation coefficient of 1.00, gives the rate of the reaction, which equals the slope of the plot. Experiments were carried out in triplicate with different acetone and acid concentrations to find the values of the other reaction orders: a and c. Results are summarized in Table 1. Volume 69 Number 7 July 1992
585
Table 2. Concentrations, Rates, and Rate Constants for the Bmmination of Acetone Measured with a Cary 14 Spectrophotometerat 25 'C
Experiment no.
[CaH60] M
[HI]
[Brzk -A[Brz]/At M
Figure 2. Iodine concentration plotted as a function of time in the iodination of acetone. Commercial Spectrophotometer
The second method for determining the rate law used bromine instead of iodine, and a Cary Model 14 recording spectmphotometer was used to measure absorbance as a function of time. Absorbance measurements were made in l-cm optical cuvettes at 395 nm, which is an absorbance maximum for aqueous bromine. In all other respects, the experiments were very similar to those carried out with the Blocktmnic. Results were also quite similar, as shown in Table 2. Visual Determination
The third method, also using bromine, involved the visual determination of the time required for complete reaction of bromine to determine rates. Here, it is not possible to determine the concentration of bromine as a function of time, so it is necessary to tell the students that the reaction is zem-order in bromine concentration. The students can then calculate rates from the times required for complete disappearance of a known initial concentration of aqueous bromine. Experimental
In this experiment ( I , 2).two test tubes are wrapped in white paper; and both are placed on a sheet of white paper. One test tube. which contains a reaction solution that has already gone'to completion, acts as a standard for color matching. The other test tube contains a reacting solution. Table 1. Concentrations, Rates, and Rate Constants for the lodination of Acetone Measured with a Blocktronic at 23 'C
Experimen1 no.
[CaHeO] M
[Hi] M
[Iz] x
lo3
M
Rate rate x lo6 -A[I~]/A~ constant ~h kx105
1
1.60
0.403
4.14
2
0.80
0.1.1
3.69
2.85
3.52
3
0.40
0.202
3.96
2.94
3.65
4
0.80
0.403
4.26
12.9
4.00
5
1.60
0.202
4.38
12.7
3.93
6
0.80
0.202
4.28
Journal of Chemical Education
28.0
5.99
4.35
3.70
+
16" = (3.86 0.2) x 10" M-'s-'
The students look down the length of both test tubes and measure the time required for all the bromine wlor to disappear, that is, the time required for both solutions to appear to have the same color intensity. Values of
are calculated by dividing the initial bromine concentration by the time required for complete reaction. The rates calculated in this way are treated mathematically the same as in the other two methods. Some typical results are shown in Table 3. Additional results obtained by one of the authors fmm students at the University of Pennsylvania were similar. The rate constant determined by a section of students averaged (3.96 0.58) x 10dM-'s" with values in the range 0f(2.60 lo4, 5.75 104). This relatively wide range probably reflects the sensitivity with which different students eould see the pale yellow of dilute bromine solutions. All three methods yield similar results for the orders a and c and for the rate constant. It is not clear why the average value fork found for the bromine reactions is higher than that found for the iodine reaction, although the slightly different temperature may be partially responsible.
+
Table 3. Concentrations, Rates, and Rate Constants for the Bromination of Acetone Measured Visually at 25 'C
Experirn [CaHsO [Ht] M ent no. ] M
[Brz] M x103
586
Rate
l o 6 Constant MIS k x lo5
x
x 10
M
time
-A[BrzgAt
s
x 10
Rate Constant
MIS
k x105
and obtain a straight line with a slope of 1.10. Again, we round this value to 1for the reaction order a. The rate law thus has the following form. -A[IzI -At - k[Cs&OI[HCI
which agrees with the accepted rate law (3).This equation can be solved to find a value of k for each experiment. These values appear in Table 1. The average value of k for all the experiments is (3.49 + 0.20) lo6 IVPS?
Figure 3. A log-log plot used to find the order c in the rate law for iodination of acetone. Determining the Rate Law The values for the reaction orders a and c were found from log-log plots (6). When the order in iodine (b) is 0, the rate law simplifies to the following expression.
Taking the log of each side of this equation results in the following equation.
In Table 2 the concentration of the acetone remains constant in experiments 1-4. For these experiments, the loglog equation simplifies to the following expression.
E]
log -- =clog ([HCI)= C where C = log k A plot of
+ a log ([C&OI).
Comparison of Techniques As a teaching tool, the experiments that use the spectrophotometer or the Blocktmnic colorimeter are more desirable than the visual method. Both instruments are able to measure the absorbance of a reacting solution over time, allowing the student to see that the reaction proceeds a t a constant rate with respect to the halogen concentration. Thus, they can determine that the order in halogen concentration b is 0. This is not possible with the visual method. A second drawback of the visual method is that students must continuallv observe a test tube that contains the reacting solution"unti1 it seems that the color has disappeared. This judgment is often difficult to make reliably and consistently. In contrast, the instrumental methods do not call for such judgment, and they give reproducible results. The Blocktronic colorimeter offers several advantages over the spectrophotometer. Cost is certainly a major advantage. The Blocktmnic calorimeter can be built for less than five dollars, while even a low-end spectrophotometer costs several thousand dollars. Even including the cost of an Apple computer, which can be put to other uses, the Blocktronic offers several economic advantages.
With other inexpensive devices, the computer can be used to measure temperature, resistance, pH, and pressure. The data collected by the computer can also be imported into a spreadsheet, and calculations can then be carried out quickly and reliably. Using a computer for data acquisition and data processing gives the students exposure to modem reseamb methods in the beginning laboratory courses. Literature Cied
should give a straight line with a slope of c. As shown in Figure 3, a straight line with a slope of 1.07 results. Since reaction orders are normally whole numbers or simple fractions, we assign a value of 1to the order c. Applying the same procedure to experiments 4-7, we plot log [*]vs.
log ([C3H@l)
mid, E. S. Mechanism and Strueturn in OIganlc Chomkty; Holt: NewYork, 1959;
- "?* p".e.
4. Adams. T;Hart-, K; Jomeay, D.;Miles, P.; Moore, J. W.I~arttruetionaforBuilding the Bloekfmnk I. LMW2: Project SERAPHIM, Univeraify ofwisemain: Madison.
WI 53706.
6. Rich, J. S.: Hartman, K.A.;Rsrmuasm,M.;Halris,M.; Babe, M. Did. 12W;Projeb SERAPHIM, University ofWiseonsim Madison,WI 53706. 6 . Birk, J. P J. Chem. Edw. 1976,53,704-107.
Volume 69 Number 7 July 1992
587
