In the Classroom edited by
JCE DigiDemos: Tested Demonstrations
Ed Vitz Kutztown University Kutztown, PA 19530
A Kinetics Demonstration Involving a Green–Red–Green Color Change Resulting from a Large-Amplitude pH Oscillation submitted by:
Brian W. Pfennig* and Richard T. Roberts Department of Chemistry, Ursinus College, Collegeville, PA 19426
checked by:
Dean Campbell Department of Chemistry, Bradley University, Peoria, IL 61625-0208
The goal of this article is to provide chemistry teachers with a chemical demonstration of a clock reaction for the winter holiday season—one that changes in color from green to red to green again. Demonstrations involving clock reactions have been used for over 50 years as a didactic tool to introduce students to many of the basic principles of kinetics. One of the most common clock reactions is the Landolt iodine clock reaction, which was first reported in this Journal in 1952 (1). The reaction involves the oxidation of iodide ion with persulfate to produce iodine and the sulfate ion: S2O82−(aq) + 2I−(aq)
I2 (aq) + 2SO42−(aq) (1)
The iodine clock reaction is pedagogically useful in the illustration of the effects of reactant concentrations and temperature on the rate of a chemical reaction, the balancing of oxidation–reduction equations, and the determination of reaction order (2). One of the main drawbacks of the demonstration is that the solutions must be prepared immediately before the presentation or the iodide will oxidize in the presence of oxygen. A quick survey of the literature, however, reveals that this problem was addressed years ago by Kauffman (3) and more recently by Mitchell (4), allowing for the preparation of materials with a much longer shelf-life. An entertaining variation on the iodine clock reaction, known as the “Old Nassau” reaction, was popularized by Hubert Alyea in the 1960s (5). Named after the oldest building on Princeton’s campus, Nassau Hall, the demonstration was timed to undergo a color change mimicking the University’s colors (from orange to black) at the precise moment when Professor Alyea completed the chorus of Princeton’s venerable anthem, “Old Nassau”. The orange coloration results from the formation of a suspension of HgI2, while the blue–black color appears when all the iodate has been consumed and the excess iodide reacts with starch to form a polyiodide ion (6): 3HSO3−(aq) + IO3−(aq) 3SO42−(aq) + I−(aq) + 3H+(aq)
+ Hg2 (aq) + 2I−(aq)
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HgI2(s) orange
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(2a)
(2b)
IO3−(aq) + 5I−(aq) + 6H+(aq) 3I2(aq) + 3H2O(l) I2(aq) + I −(aq) + starch
(2c)
blue–black polyiodide (2d) ion on starch
Inspired by the orange–black changes of the “Old Nassau” or Halloween reaction, our goal was to develop a chemical demonstration of a clock reaction for the winter holiday season, namely one that changed color from green to red. One of the easiest ways to achieve this color change was to use different acid–base indicators, along with a clock reaction having large-amplitude pH changes. One such example is the formaldehyde–sulfite clock reaction described in Shakhashiri’s volumes on chemical demonstrations (7).1 Somewhat surprisingly, however, there were very few systems in the chemical literature that showed both large-amplitude and periodic oscillations in pH (8). The most promising candidate for a periodic pH demonstration was a series of reactions developed by Rabai and Beck in 1988 (9), known as the IST system (for iodate, sulfite, and thiosulfate) (10). The key reactions2 are SO32−(aq) + H+(aq)
HSO3−(aq)
IO3−(aq) + 3HSO3−(aq) I−(aq) + 3SO42−(aq) + 3H+(aq)
(3a)
(3b)
− − + IO3 (aq) + 6S2O32 (aq) + 6H (aq) (3c) − − I (aq) + 3S4O62 (aq) + 3H2O(l)
In order for the pH to oscillate, there must be two competing processes. These processes are the reduction of iodate ion by either bisulfite or thiosulfate. At high pH (pH > 4), the iodate is reduced by the bisulfite ion to make iodide and sulfate, according to eq 3b. Because sulfate is a weaker base than sulfite, there is a concomitant release of protons. However, the protons are in equilibrium with sulfite ion and bisulfite, as shown in eq 3a, so that the concentration of H+ increases only slightly until all the sulfite is converted. At this point, the pH rapidly decreases until a minimum is reached
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as the bisulfite is consumed. At low pH, the iodate ion is reduced in the presence of acid by thiosulfate, according to eq 3c. This latter reaction consumes hydrogen ions, so that the pH begins to increase again. Under a very narrow range of initial concentrations3 the pH will oscillate through more than one cycle (10). For purposes of the demonstration, a combination of three acid–base indicators was employed: bromocresol green, bromothymol blue, and methyl yellow. Under the initial conditions of the experiment (pH ∼6.5), bromocresol green and bromothymol blue are both blue in color, while methyl yellow is yellow. The combination of blue and yellow gives the solution a green appearance. The use of both bromocresol green and bromothymol blue ensures that the solution will remain sufficiently green in color until the pH has dropped by several pH units. At pH < 3, both bromocresol green and bromothymol blue are yellow in color, while methyl yellow is red. Because the red color is more intense than the yellow, the reaction mixture appears red to red–orange at low pH. As the thiosulfate is consumed, the protons are consumed in eq 3c and the pH rises rapidly. Under a very narrow range of experimental conditions (10), the system can cycle through several oscillations, each of which is significantly smaller than the former. The first several damped oscillations can be detected if the pH is recorded as a function of time (9). Because the pH changes decrease with each oscillation, only the first oscillation can be observed using the acid–base indicators in this chemical demonstration. We have optimized the experimental conditions from those in the literature (10) so that the initial pH spike is maximized. This ensures that the color change of the acid–base indicators will be observed during the demonstration. Experimental Section The chemical demonstration was performed in a 250mL beaker equipped with a magnetic stir bar. The total volume of the solutions after mixing was 100.0 mL. The following stock solutions were prepared in advance: 0.0600 M KIO3 (0.321 g兾25.0 mL), 0.150 M Na2SO3 (0.378 g兾20.0 mL), and 0.100 M Na2S2O3 (0.237 g兾15.0 mL). A 0.0400 M H2SO4 stock solution was made by diluting 27.8 mL of concentrated sulfuric acid in a 500.0-mL volumetric flask and then further diluting 10.0 mL of this solution into a second 500.0-mL volumetric flask. [Caution: concentrated sulfuric acid can cause severe burns; in order to avoid spattering, add the acid to the water when doing the dilution]. All of the starting materials were purchased from Sigma-Aldrich and were of analytical grade. Deionized, distilled water was used for all of the dilutions. Because sulfite can undergo air-oxidation, the sulfite solution was prepared just prior to performing the chemical demonstration. The indicator solutions were prepared as follows. Bromocresol green was made according to the literature procedure (11) by dissolving 0.1 g of the solid in 14.3 mL of 0.01 M NaOH and diluting with 253.7 mL of water. The bromothymol blue was a 0.1% (by mass) solution in 50:50 (v:v) ethanol:water (11). The methyl yellow solution was 0.010 M and was prepared by dissolving 0.0358 g of methyl yellow solid in 15.6 mL of methanol, which yields a concentration of 0.0102 M.
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To perform the demonstration, the sulfite stock solution (20.0 mL) was placed into a 250-mL beaker equipped with a magnetic stirring bar. Five drops of bromocresol green, five drops of bromothymol blue, and ten drops of methyl yellow were then added. Next, 37.0 mL of the 0.0400 M sulfuric acid solution and 2.0 mL of water were added, followed by 15.0 mL of the stock thiosulfate solution. To initiate the oscillatory reaction, the 25.0 mL stock solution of iodate ion was added last. The initial color of the solution was green. The pH of the reaction mixture was monitored as a function of time using a model RJ combination pH probe from Analytical Sensors, Inc. connected to a Dell Latitude C540 laptop computer using a LabWorks II station interface. The pH meter was first calibrated using pH 7.00 and 4.01 buffer solutions from Sigma-Aldrich. The color changes of the acid– base indicators from green to red to green again were simultaneously recorded using a standard digital camera in video mode. For larger classroom settings, the use of a Webcam with computer projection is recommended. A general explanation of kinetics and the specific reactions involved in the IST system was discussed while the class waited for the color changes to occur. Discussion Immediately after mixing, the pH of the green solution started around pH = 7, gradually dipped to pH < 4 (where the indicator is red) and then rebounded back above a pH > 6.5 (green). The red color persisted only briefly (as a clearly perceptible flash) before the pH rose again and the indicator color changed back to green. The length of time until the appearance of the green-to-red color change was recorded. The demonstration was performed by a senior chemistry major specializing in secondary education on various occasions to a variety of audiences throughout the academic year. Oscillatory (spike) behavior was observed on all but one of these attempts (on the failed occasion, the pH gradually decreased, the indicator changed from green to red but it never rebounded). On average, the indicator changed color from green to red approximately 4.5 ± 2.0 min after mixing, stayed red for about 1 s and then switched back to green. A typical plot of the data is shown in Figure 1. The conditions used for the chemical demonstration were optimized so as to enhance the magnitude of the initial pH spike, so that only a single oscillation was observed using these initial concentrations. While the reaction is proceeding, the instructor can present a brief discussion of the conditions necessary for oscillatory kinetics. Lotka (12) was the first to postulate a general mechanism for oscillating reactions. His original model is presented in eq 4. Two independent processes are occurring with a third reaction that links the two. The system is autocatalytic in X:
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A
Y
(4a)
X
P
(4b)
X + Y
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2X
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(4c)
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In the Classroom
Figure 1. Clock behavior in the IST demonstration showing pH as a function of time after mixing and the colors of the indicator solution before, at, and following the spike in pH. The dark data points indicate the green-colored solution and the gray indicate the redcolored solution. Conditions: [IO3−] 0 = 0.0150 M, [SO32−] 0 = 0.0300 M, [S2O32−]0 = 0.0150 M, [H2SO4]0 = 0.0148 M, 5 drops bromocresol green, 5 drops bromothymol blue, 10 drops methyl yellow, 100-mL total volume, T = 23 ⬚C.
The original Lotka model is overly simplistic, however, because it assumes that the elementary steps are not accompanied by more complex reactions in the overall mechanism. A modified form that accounts for higher-order steps is A + B
A + B + X A + X + Y
Notes 1. The checker reports that Flinn Scientific, Inc. supplies a demonstration kit (Catalog No. AP6858) based on the formaldehyde clock, for a reaction that undergoes a single change from green to red. 2. The reactions are not implied to be elementary steps in a mechanism. 3. The initial concentrations must meet the following criteria: • [IO3−]0兾{([HSO3−]0兾3) + ([S2O32−]0兾6)} is slightly larger than one, and • [S2O32−]0 falls between 2.1–2.2 [H2SO4]0.
(5a)
Y
P1
(5b)
2X + P2
(5c)
In the context of the IST system used in this demonstration, the X, Y, A, B, P1, and P2 roles correlate to H+, HSO3−, IO3−, S2O32−, S4O62−, and SO42−, respectively. Because clock behavior will only be observed over a narrow range of concentrations, Rabai and Beck (9) have derived an empirical set of requirements for oscillatory behavior to be observed. Using initial concentrations, the following sets of inequalities must be employed: (a) 3[IO3−] < ([SO32−] + [S2O32−]), (b) 3[IO3−] > [SO32−], and (c) [H+] > [S2O32−].
The demonstration is appropriate for high school or introductory-level college chemistry courses because it does not require highly specialized or expensive materials, the reaction
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Literature Cited 1. 2. 3. 4. 5. 6. 7.
8. 9. 10. 11.
Conclusion
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occurs in a reasonable period of time, and there are no real safety considerations. In addition to serving as an entertaining demonstration for the winter holiday season, the IST clock reaction presented in this publication provides a springboard for the discussion of the factors that influence the rates of chemical reactions. One advantage that this demonstration has over the similar formaldehyde–sulfite reaction in the literature (7) is that its periodic (spike) behavior can also be used to illustrate the requirements for oscillatory kinetic behavior. Furthermore, the demonstration presents a fine example of how sulfate is a weaker base than sulfite. The negative charges on the sulfate ion are stabilized by resonance to a greater extent than those on the sulfite ion. Thus, the sulfate ion will be less likely to attract a proton than the sulfite ion will be. The demonstration can also be used to review the concept of oxidation numbers and the balancing of oxidation–reduction reactions.
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12.
Evans, G. G. J. Chem. Educ. 1952, 29, 139. Creary, X.; Morris, K. M. J. Chem. Educ. 1999, 76, 530. Kauffman, G. B.; Hall, C. R. J. Chem. Educ. 1958, 35, 577. Mitchell, R. S. J. Chem. Educ. 1996, 73, 783. Alyea, H. N. J. Chem. Educ. 1965, 42, 19. Fortman, J. J.; Binford, J. J. Chem. Educ. 1992, 69, 236. Shakhashiri, B. Z. Chemical Demonstrations: A Handbook for Teachers of Chemistry; University of Wisconsin Press: Madison, WI, 1985; Vol. 4, pp 70–74. Rabai, G.; Orban, M.; Epstein, I. R. Acc. Chem. Res. 1990, 23, 258. Rabai, G.; Beck, M. T. J. Phys. Chem. 1988, 92, 2804. Rabai, G.; Beck, M. T. J. Phys. Chem. 1988, 92, 4831. Gordon, A. J.; Ford, R. A. The Chemist’s Companion: A Handbook of Practical Data, Techniques, and References; John Wiley & Sons: New York, 1972. Lotka, A. J. J. Phys. Chem. 1910, 14, 271.
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