In the Classroom edited by
Overhead Projector Demonstrations
Doris K. Kolb Bradley University Peoria, IL 61625
A New Twist on the Iodine Clock Reaction: Determining the Order of a Reaction Xavier Creary and Karen M. Morris* Department of Chemistry and Biochemistry, University of Notre Dame, Notre Dame, IN 46556
Over the years chemical educators have investigated, discussed, and used clock reactions in their classrooms to illustrate chemical kinetics (1–3). Other chemical educators have experimented with the Landolt iodine clock reaction, and their results have often been published in this Journal (4, 5). The Landolt iodine clock reaction is a reliable and well-used chemistry demonstration owing to a variety of features: ease of solution preparation, striking color change indicating reaction completion, convenience in changing reaction “clock time”, and effective, dramatic presentation. The iodine clock reaction can also be used to illustrate the kinetic order of a reaction, and an overhead projector demonstration was developed three years ago for General Chemistry classes at the University of Notre Dame showing this concept. This demonstration has been used successfully with consistent results since that time. Demonstration Preparation The solution preparation for this demonstration is slightly different from other preparations (refer to 1, 3, 4 ) because sodium metabisulfite (Na2S2O5) is used instead of sodium bisulfite (NaHSO3). Metabisulfite serves as a source of HSO3{ according to the equation S2O52{(aq) + H2O(,) → HSO3{(aq) Since the iodine clock reaction depends on the reducing properties of the bisulfite ion, using sodium metabisulfite is equivalent to the use of sodium bisulfite. If possible, the prepared solutions should be used promptly, since solutions containing bisulfite ions are subject to air oxidation. Additionally, although the bisulfite solution is dilute, gaseous sulfur dioxide can also be evolved, since sulfuric acid is added to this solution (1). Keeping the bisulfite solution storage flask tightly capped will reduce exposure to atmospheric oxygen and reduce loss of gaseous SO2. To prepare solution A, dissolve 4.30 g of KIO3 in 200 mL of water in a 1-L volumetric flask. Add enough water to the flask to make 1 L of solution. Stopper and label the flask. To prepare solution B, heat 300 mL of water to boiling in a 1-L beaker (a hot-plate–stirrer works well). Dissolve 4.0 g of soluble starch in the water; the resulting solution should be translucent. Dilute the starch solution with water to a
volume of 900 mL and allow it to cool. Dissolve 0.20 g of sodium metabisulfite (Na2S2O5) in the cooled starch solution. Pour the solution into a 1-L volumetric flask. Add 0.49 g of concentrated sulfuric acid (H2SO4) to the solution in the flask, stopper the flask, and mix the contents. Add enough water to the flask to make 1 L of solution, tightly stopper the flask to reduce exposure to atmospheric oxygen (use stopcock grease if necessary), and mix the contents. Label flask. Demonstration Setup Obtain four glass, 90 × 50-mm crystallizing dishes (250mL beakers may also be used; however the crystallizing dishes work best on the overhead projector). Label the dishes on their sides, consecutively, from one to four. Using a 25-mL graduated cylinder, measure the amount of solution A listed in Table 1 for dish 1 and add enough water to make the total solution volume 25 mL. Repeat for the other dishes using the volumes of solution A and water in Table 1. Just before the demonstration, fill four 25-mL graduated cylinders with solution B to the 25-mL mark. Place the four dishes containing the different concentrations of solution A and the graduated cylinders containing solution B near an overhead projector. Place a glass stirring rod nearby as well. A timer or stopwatch capable of timing to tenths of seconds will also be necessary to do this demonstration. Doing the Demonstration Place the first dish containing solution A on the overhead projector and turn on the projector so that the image is displayed. Very quickly add solution B to solution A on the overhead and stir briefly to make a homogenous solution. Start timing the reaction as soon as solution B is added. Record the time it takes for the solution to turn completely blue. Repeat this process for the remaining dishes. After the reaction times have been collected, the data must be converted to reaction rates because a plot of ln rate vs ln [IO 3{] is needed to determine the reaction order (see Discussion section). The method of initial rates, where rate = ∆[HSO3{]/∆t, is used because the [HSO3{] is very small and constant (10 {4 M). HSO3{ is also assumed to react completely (6 ). The rate values, therefore, are determined for each reaction in which rate = 10 {4 mol L{1/reaction time.
*Email:
[email protected].
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Journal of Chemical Education • Vol. 76 No. 4 April 1999 • JChemEd.chem.wisc.edu
In the Classroom
Disposal and Safety
Table 1. Preparation of Solution A
Since the bisulfite solution is acidic, some gaseous sulfur dioxide can be released, which can irritate the respiratory system. CAUTION: People hypersensitive to sulfites should avoid direct contact with the solutions in this demonstration (1). To dispose of the products of this demonstration, combine the deep blue solutions in a large beaker. While stirring the mixture, slowly add solid sodium thiosulfate (Na2S2O3?5H2O) until the mixture is no longer blue. Flush the mixture down the drain with large quantities of water (1). Discussion
Solution A Vol/mL
Dish No.
Water Vol/mL
[IO3 ]
{
1
5
20
0.004
2
10
15
0.008
3
15
10
0.012
4
20
5
0.016
Table 2. Data from General Chemistr y, Spring 1997 {
The concentration of the bisulfite solution is very small when compared to the iodate solution ([bisulfite] ~10{4 M; [IO3{] = 0.004–0.016 M). Since early studies of the Landolt reaction indicated that the reaction is first order in [IO3{] and first order in [HSO 3{] (1) and since the [IO3{] doesn’t change appreciably during the course of the reaction, the rate expression is pseudo-first-order and can be written as rate = kobs[IO3{]1 (6 ). Since students do not know the reaction order, the rate expression for the reaction is written as rate = kobs[IO3{]n. Taking the natural logarithm of both sides gives ln rate = ln(kobs[IO3{]n), which is converted to ln rate = n ln[IO3{] + ln kobs, a straightline equation with slope = n (order of reaction). Using classroom data, therefore, the order of reaction can be determined by plotting ln rate vs ln[IO3{]. Refer to the graphing example below (Table 2, Fig. 1).
[IO3 ]/ mol L{1
{
[HSO3 ]/ Time/ mol L{1 s
Rate/ { Ln [IO3 ] (mol L{1 s {1 )
Ln Rate
10
{4
0.008
10
{4
26
10 /26
{4.828
{12.468
0.012
10{4
18
10{4/18
{4.423
{12.101
0.016
10{4
14
10{4/14
{4.135
{11.849
0.004
56
{4
{5.521
{13.236
{4
10 /56
Literature Cited 1. Shakhashiri, B. Chemical Demonstrations; The University of Wisconsin Press: Madison, WI, 1992; Vol. 4, pp 3–25. 2. Lambert, J. L.; Fina, G. T. J. Chem. Educ. 1984, 61, 1037–1038. 3. Tested Demonstrations in Chemistry; Alyea, H. N.; Dutton, F. B., Eds.; Journal of Chemical Education: Easton, PA, 1965; p 19. 4. Autuoir, M. A.; Brolo, A. G.; Mateus, Al. L. J. Chem. Educ. 1989, 66, 852–853. 5. Brice, L. K. J. Chem. Educ. 1980, 57, 152. 6. Bromberg, J. P. Physical Chemistry, 2nd ed.; Allyn and Bacon: Newton, MA, 1984; pp 890–892.
Figure 1. Order of reaction for the iodine clock. Slope = n =1.00; r = .998.
JChemEd.chem.wisc.edu • Vol. 76 No. 4 April 1999 • Journal of Chemical Education
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