Clock Reaction Revisited: Catalyzed Redox Substrate-Depletive

Communication pubs.acs.org/jchemeduc

Cite This: J. Chem. Educ. XXXX, XXX, XXX−XXX

Clock Reaction Revisited: Catalyzed Redox Substrate-Depletive Reactions Taweetham Limpanuparb,* Chattarin Ruchawapol, and Dulyarat Sathainthammanee Mahidol University International College, Mahidol University, Salaya, Phutthamonthon, Nakhon Pathom 73170, Thailand

Downloaded via TULANE UNIV on February 14, 2019 at 23:35:28 (UTC). See https://pubs.acs.org/sharingguidelines for options on how to legitimately share published articles.

S Supporting Information *

ABSTRACT: Quantitative demonstration of chemical kinetics often requires sophisticated equipment and materials. By using inexpensive and commonly available chemicals, we propose here two clock reactions, oxygen−safranin−benzoin clock and cysteine−iodine− hydrogen peroxide clock. Students simply measure and add stock solutions to start clock reactions. Sudden discoloration or coloration of the reaction mixture marks the end of the reaction, and this can be conveniently timed. Students then complete calculations and graphing in a spreadsheet file by following common protocols. Rates of reaction are calculated from the ratio of [limiting reactant]0 and clock time. Orders of reactions and rate constants are determined by the method of initial rate. Additional discussions on topics such as catalysis, reaction mechanism, redox reaction, and pseudo-zeroth-order reaction are also possible. KEYWORDS: First-Year Undergraduate/General, Physical Chemistry, Hands-On Learning/Manipulatives, Computer-Based Learning, Oxidation/Reduction, Kinetics, Rate Law, Catalysis, Laboratory Computing/Interfacing

■

A number of papers19−22 discussed timing the discoloration of the blue color of methylene blue implying a clock behavior in these experiments. A direct link between the blue bottle reaction and clock reaction has also been suggested in a theoretical paper by Limpanuparb et al.16 In this communication, a rapid version of the blue bottle reaction14 is used for further investigation. Experimental confirmation that the blue bottle reaction has the clock behavior is provided in the next section.

CLOCK REACTIONS There are three seminal reviews on clock reactions by Schreiner et al.1 in 1992, Lente et al.2 in 2007, and Horvath and Nagypal3 in 2015. A clock reaction derives its name from an alarm clock. The alarm goes off by an abrupt change in the concentration of a clock species after a time lag (induction period or clock time). In the most recent review article, Horvath and Nagypal3 classified clock reactions into three main groups: substrate-depletive clock reaction, autocatalysis-driven clock reaction,4,5 and pseudoclock behavior. Several subsequent papers,6−8 however, suggest that depending on the circumstance some reactions such as the iodate−arsenous acid reaction may behave as a substratedepletive clock reaction, an autocatalysis-driven clock reaction, or a crazy clock reaction.

Iodine Clock

There are a number of variations of the clock reaction involving iodine species. The oldest known clock reaction is the 1886 Landolt iodine clock23,24 where the colorless solution made of iodate, bisulfite, and starch suddenly turns dark blue after an induction period. The old Nassau clock reaction is similar to the Landolt but involves a precipitation of mercury(II) iodide.1 The most widely used clock reaction in an educational setting is possibly the oxidation of a reducing agent catalyzed by iodine.1,25−42 The preparation of reducing agent−iodine− hydrogen peroxide is relatively safe and simple, and the mechanism is well-understood. Various modifications of reducing agents, e.g., cysteine16 and vitamin C,31,32 for this reaction have also been proposed. Following the most recent modification of the iodine clock,16 cysteine−iodine−hydrogen peroxide is the basis for further investigation in this communication.

Blue Bottle Reaction

The classical blue bottle reaction9,10 is made of caustic soda (NaOH or KOH), glucose, and methylene blue in water. The overall reaction is the oxidation of glucose by dissolved oxygen into products11 catalyzed by a redox indicator. Methylene blue oxidizes glucose, and after a while, most of it is turned into leucomethylene blue, a colorless compound. Shaking the solution brings in more atmospheric oxygen which oxidizes the leuco form of the indicator back to its original blue form. Since the rate of reaction for methylene blue and glucose is slower than that of leucomethylene blue and oxygen, the solution remains blue until oxygen is consumed.12 There are various variations of the experiment using different reducing agents (e.g., vitamin C, benzoin, and acetoin13−16) and dyes (e.g., safranine, indigo carmine, and erioglaucine13−15,17−19). © XXXX American Chemical Society and Division of Chemical Education, Inc.

Received: July 11, 2018 Revised: January 29, 2019

A

DOI: 10.1021/acs.jchemed.8b00547 J. Chem. Educ. XXXX, XXX, XXX−XXX

Journal of Chemical Education

Communication

Table 1. Stock Solution Components for the Clock Reactions Clock Reaction Oxygen−safranine−benzoin Cysteine−iodine−hydrogen peroxide

Beaker A Water, benzoin (C14H12O2), safranin (C20H19ClN4) Water, starch, buffer, cysteine (C3H7NO2S), potassium iodide (KI)

Theoretical Framework

and clock times, the rate of reaction, the order of reaction, and the rate constant can be calculated following standard procedures in general chemistry textbooks. We aim to use the experiments as a laboratory exercise for first-year undergraduate students of nonchemistry majors to supplement lecture in chemical kinetics. Various versions of the two reactions have been used in ICCH224 Integrated Laboratory Techniques in Chemistry I class at Mahidol University International College for a few years. In the past year, the two reactions were combined and presented to students in one 4 h laboratory session. Among many other possibilities for the clock reaction experiment, we chose these two reactions for our class because of the impressive visual effect, similarly suitable clock time, ease of experiment/interpretation, and availability/cost/toxicity of chemicals. We used benzoin and cysteine as reducing agents not only because they are recently discovered reactions but also because they represent important concepts in chemistry. The use of benzoin in the blue bottle reaction conclusively disproves the commonly accepted belief that the reaction is autoxidation of aldehyde to carboxylic acid. On the other hand, the use of cysteine, a conditionally essential amino acid, in the iodine clock reaction leads to discussion of the antioxidant property of biomolecules in class. It is important to note that while iodine may be used for wound treatment, the chemical “benzoin” discussed here is not an ingredient in tincture of benzoin used for similar medical purposes. There are two basic concepts behind these simple timing experiments.

For the purpose of this communication, only two specific systems, oxygen−safranin−benzoin (the blue bottle reaction) and cysteine−iodine−hydrogen peroxide (the iodine clock), are discussed. These two systems are catalyzed redox substratedepletive clock reactions. The names of the two reactions here are provided in the format “limiting reactant−catalyst−other reactant”. Reactants are prepared in two separate portions, and they are combined to start the reaction as shown in Table 1. During the reaction, limiting reactants (oxygen and cysteine) are consumed by other reactants (benzoin and hydrogen peroxide) at a constant rate with the aid of catalyst (safranin and iodine, respectively). After the limiting reactants are used up, catalysts can no longer be converted to their other forms, and this leads to a sudden change in color of the reaction mixture. Overall reactions for the two clock reactions according to the Supporting Information are O2 + C14 H12O2 + OH− → C14 H10O2 + HO2− + H 2O (1)

and H 2O2 + 2C3H 7NO2 S → C6H12N2O4 S2 + 2H 2O

(2)

respectively. Unlike the equilibrium constant expression, the rate law for the two clock reactions may involve not only reactants but also catalysts. In other words, all chemical reagents, in theory, may appear in rate laws rclock1 = kclock1[O2 ]g [C14H12O2 ]h [OH−]i [C20H19N4 +] j

(3)

rclock2 = kclock2[H 2O2 ]l [C3H 7NO2 S]m [I−]n [H3O+] p

(4)

• First, most chemical reactions, in theory, have no completion time. The rate of reaction becomes slower as concentrations of reactions decrease, and hence, mathematically it requires infinite time to complete these reactions. The only exception is a zeroth-order reaction where the rate of reaction is constant. Clock reactions have a finite reaction time due to this zerothorder reaction behavior.

where r stands for the rate of reaction. For simplicity, concentrations of two chemicals are varied in the first clock and three chemicals in the second clock are varied. (To vary the concentration of oxygen and dye, one would require more complicated machines for measurement, and to vary the concentration of H3O+ one would need to play with a different buffer solution.) The rate laws for our data analysis are therefore rclock1 = kobs1[C14H12O2 ]h [OH−]i

• Second, the observed order of reaction generally appears to be the order of the reaction with respect to limiting reactant if the concentration of the limiting reactant is much lower than those of other reactants. This phenomenon is called a pseudo-order reaction. In particular, the observed order of reaction may appear to be zeroth-order if the reaction is zeroth-order with respect to its limiting reactant. Experimental conditions for two clock systems in this communication are carefully set-up so that limiting reactants are present in a relatively small amount. The order of reaction with respect to limiting reactants is zeroth because their reactions with catalysts are fast and hence not the rate-determining step.

(5)

and rclock2 = kobs2[H 2O2 ]l [C3H 7NO2 S]m [I−]n

(6)

where the order of reaction, h, i, l, m, and n are determined and rounded to integer values first and the two observed rate constants kobs1 and kobs2 are determined from the rate law. See reaction schemes in Supporting Information for more details.

■

Beaker B Sodium hydroxide (NaOH) Hydrogen peroxide (H2O2)

JUSTIFICATION

Pedagogical Applications

Students appreciate these concepts by varying concentrations of reagents and observing changes in the reaction time. The clock time increases linearly with respect to the concentration of the limiting reactant. The clock time may decrease if the concentrations of other reagents are increased. The behavior of clock time vs initial concentration is used to derive/confirm the order of reaction and the rate law.

Quantitative chemical kinetic experiments often require sophisticated equipment to monitor the concentration of involved chemical species with respect to time.43,44 Most clock reaction experiments, on the other hand, only require students to measure liquid solutions, mix them, and time the experiment from the mixing to the color change. With initial concentrations B



Communication

Figure 1. Graphs of absorbance and dissolved oxygen vs time. Left: Absorbance at λ = 518 nm and dissolved oxygen values vs time in the oxygen− safranine−benzoin clock reaction. Right: Absorbance at λ = 575 nm vs time in the cysteine−iodine−hydrogen peroxide clock reaction.

General chemistry textbooks45,46 mainly discuss the pseudofirst-order reaction and may extend it to the pseudo-secondorder reaction only. As a result, the pseudo-zeroth-order reaction may appear to be unfamiliar or impossible to many students. Chemical clock reactions here and many other biological clocks are tangible examples where a pseudo-order reaction can be of zeroth order. Unlike negative pH in which the acid must be strong and concentrated,47 there is no additional experimental difficulty to explore a pseudo-zeroth-order reaction. Students are encouraged to connect the misconception regarding nonexistence of negative pH and nonexistence of pseudo-zeroth-order reaction to other areas of unexplored chemistry and science.

discussion or to convince them that it is possible to time these reactions accurately by bare eyes. (If experiments are conducted by the same student under the same experimental conditions, clock times are reproducible with no more than a few seconds discrepancy.)

■

EXPERIMENTAL SECTION The aim of these experiments is to study the effect of initial concentration of each reagent on the clock time. Experimental work may be completed in 2 h. During our 4 h lab session, the first hour is for lecture and demonstration and the last hour is for data analysis and writing up. Students worked in a group of two or three so that they can take turns to time the experiment and prepare a new batch of solutions. The laboratory manual in the Supporting Information contains the recommended concentrations of stock solutions, the equipment used for measurement, and the proportion of each stock solution for the experiments. Supporting Information files are intended for instructor use. The only documents provided to students are the laboratory manual and spreadsheet file (only “benzoin” and “cysteine” sheets). The parts of the documents that we leave blank for students to complete are highlighted in gray.

Confirmation of Clock Behavior

To confirm the clock behavior of the two reactions, we measured the absorbance of the reaction mixtures over a period of time by using the SPECTRONIC 200 spectrophotometer. For the oxygen−safranine−benzoin clock, the EXTECH SDL150 dissolved oxygen (DO) probe is used in addition to the spectrophotometer. The reaction mixture was initially prepared in a beaker with DO probe and then drawn from the beaker for spectrophotometric measurement so that the absorbance profile and oxygen profile can be matched. For accurate data acquisition by DO probe and spectrophotometer, we have intentionally made the solution of the oxygen−safranin−benzoin clock in the video take a longer time than the actual experiments that students are expected to conduct. (This is to allow time for us to set-up the probe and spectrophotometer. This also helps us obtain better data points as the sampling rate and response time of the two pieces of equipment are limited. The actual experiments for students have more sudden discoloration.) Figure 1 shows the abrupt change in the absorbance of the two clock reactions at the time for the color to change. The left plot of Figure 1A shows that DO drops almost linearly from ∼8 to ∼0 mg/L. When oxygen, which is the limiting reactant, is used up the absorbance suddenly plummets from its stable value. Our experiment is similar to a popular online video,48 but our results clearly suggest a different mechanism where color starts fading only after the oxygen is depleted. According to the right plot of Figure 1, an abrupt increase in absorbance is associated with the change of color of the solution from colorless to dark blue. Video recordings showing the visual behavior and readings from spectrophotometer/DO meter for both reactions are available in the Supporting Information. This confirmation experiment is not intended to be part of student experiment work. It can be done as a class demonstration, or results can be shown to students to facilitate

Materials and Methods

For the oxygen−safranine−benzoin clock, there are three stock solutions, benzoin, safranine, and sodium hydroxide. For the cysteine−iodine−hydrogen peroxide clock, there are five stock solutions, acetate buffer, soluble starch, L-cysteine, potassium iodide, and hydrogen peroxide. Unless otherwise specified below, all of the stock solutions are freshly prepared in distilled water at room temperature. All chemicals are used as received from suppliers without further purification. • To increase solubility, benzoin is dissolved in pure methanol. • Cysteine is dissolved quickly with a minimal amount of fuming hydrochloric acid (HCl). (It is added dropwise until cysteine is completely dissolve before adding distilled water.) • Soluble starch is stirred under heat until the solution become clear. Procedures for the two clock reactions are similar. Two beakers of solutions, A and B, are freshly prepared from stock solutions according to the laboratory manual. The solution from Beaker A is thoroughly mixed and quickly poured into Beaker B. The reaction mixture is stirred briefly by a glass rod. The reaction is timed at the beginning of mixing and terminated when the solution starts to fade (oxygen−safranine−benzoin C



Communication

Table 2. Summary of Results for the Two Clock Reactions from 14 Groups of Studentsa Median

Mean

Order of reaction for

Kinetic Parameter Benzoin Sodium hydroxide Cysteine Potassium iodide Hydrogen peroxide

0.70 0.73 0.03 0.77 0.64

Min

1.08 0.98 0.16 0.95 0.95

1.07 0.98 0.17 0.94 0.92

1.42 1.14 0.44 1.09 1.02

Max

0.19 0.12 0.10 0.09 0.10

SD

Second-order rate constant kobs1/M−1 s−1 for rclock1 = kobs1[C14H12O2]1[OH−]1 by varying the concentration of

Benzoin Sodium hydroxide

0.034 0.031

0.057 0.051

0.056 0.048

0.071 0.056

0.010 0.007

Second-order rate constant kobs2/M−1 s−1 for rclock2 = kobs2[H2O2]1[C3H7NO2S]0[I−]1 by varying the concentration of

Cysteineb Cysteinec Potassium iodide Hydrogen peroxide

0.0033 0.0063 0.0062 0.0047

0.0061 0.0074 0.0074 0.0071

0.0061 0.0074 0.0073 0.0071

0.0086 0.0090 0.0084 0.0083

0.0012 0.0006 0.0007 0.0009

r2 of tclock vs 1/[reagent]0 (for benzoin, sodium hydroxide, potassium iodide, hydrogen peroxide) or r2 of tclock vs 1/[cysteine]0

Benzoin Sodium hydroxide Cysteine Potassium iodide Hydrogen peroxide

0.9157 0.9385 0.9252 0.9293 0.8043

0.9767 0.9778 0.9982 0.9929 0.9958

0.9724 0.9786 0.9880 0.9856 0.9798

0.9985 0.9993 0.9997 0.9989 0.9999

0.0237 0.0195 0.0209 0.0195 0.0513

a

Rate constants are calculated from slopes of rate vs concentration with an exception of cysteine. bFrom the intercept of a graph of rate vs concentration. cFrom the average rate of all experiments.

proportions of stock solutions, and the reaction time. Third, order of reaction and rate constant are determined by the method of initial rate.50 There are three plots of dependent variable vs independent variable: • log r vs log [reagent]0 to determine the order of reaction, h, i, l, m, and n; • r vs [reagent]0 to confirm the order of reaction and derive kobs from the slope; • clock time vs 1/[reagent]0 (with one exception of clock time vs [cysteine]0) also to confirm the order of reaction. We can use this to prepare a reaction with any desired clock time. To avoid confusion, it is important to note the following: • Clock times, tclock, in our experiment are the total times to complete a pseudo-zeroth-order reaction for experiments with different initial concentrations, [reagent]0. This should not be confused with the concentration vs time relationship ([reagent]t vs t) commonly plotted in most textbooks. The appropriate interpretation of clock times is to convert them into reaction rates. The rate vs concentration graphs (rate vs [reagent]0) from the two clock reactions are consistent with the presentation in general chemistry textbooks. • Although the reaction rate is obtained by a pseudo-zerothorder reaction model (in the second step above), the reaction rates can be used to find actual order of reaction with respect to each reagent (in the third step above). • Unlike most clock experiments where two different concentrations per reagent are studied and experiments are done in triplicate to average out error, our experiment has four concentration points per reagent. If an outlier point is seen on a graph, it may be excluded.

clock) or turn dark (cysteine−iodine−hydrogen peroxide clock). The 50 mL reaction mixture is then discarded in a designated waste container. Beakers and glass rods are washed with tap water and rinsed with deionized (DI) water before the next batch of experiments. There are four data points each for the variation of initial concentrations of benzoin, sodium hydroxide, cysteine, potassium iodide, and hydrogen peroxide. Since some data points are shared, in total, there are 7 and 10 experiments for the oxygen−safranine−benzoin clock and the cysteine−iodine− hydrogen peroxide clock, respectively. It is our intention to perform as little as 17 experiments but derive as much as information as possible out of the experimental data. A spreadsheet file is available for student data entry. Students can instantly see if a good trend is observed (r2 > 0.95 for clock time vs initial concentration or clock time vs 1/initial concentration). Some experiments can be seen in a graph as outliers. These can be repeated if time permits to obtain a better trend. As our clock times are in a reasonable range (from 33 to 141 s at room temperature of 25 °C), we instruct students to work on one experiment at a time rather than run parallel experiments.21,26 Data Analysis

The use of a spreadsheet program is helpful to shorten the calculation time, but it is not a substitution for proper understanding of the whole process. In our case, students fill in timing data into the instructor-prepared spreadsheet file. All results are shown to students automatically without the need to type in any mathematical formula. However, students manually create their graphs, interpret results, and write their own conclusions. There are three steps to analyze these experimental data. First, model data of rate vs concentration and concentration vs time from zeroth-, first-, and second-order reactions are used to remind students of differential and integrated rate laws from the lecture class.49 Second, students are instructed to use the pseudo-zeroth-order reaction model to find the rate of reaction for each experiment from concentrations of stock solutions,

■

HAZARDS Sodium hydroxide and hydrochloric acid are corrosive. Hydrogen peroxide is a strong oxidizing agent. Methanol is toxic, volatile, and flammable. Safranin may stain clothes and D



Communication

scores for sufficiency of their knowledge are 3.4 for the experiments and 3.1 for the calculations. Student free listing responses also indicate the following: Students enjoyed watching the color change. The experiment was interesting, beautiful, and fun to perform. It is important to note that students associated being able to see by their own eyes with their understanding of the reactions. Students appreciated the sensitivity of the rate of reaction to composition of the reaction mixture. They also practiced and reinforced various techniques to properly measure volumes of stock solutions during the experiment. Spreadsheet software is helpful to simplify the process and save time for data analysis. Students can quickly see how to derive the order of reaction and rate constants from experimental data with the use of the software. However, some students did have a valid concern that without proper explanation the use of a computerized worksheet may hinder student learning. This view is reflected in relatively lower average rating scores for the two calculation activities (order of reaction and rate constant in Figure 2).

glassware. Buffer solution has a strong vinegar odor. Students are expected to work with small amounts of prepared stock solutions that generally have lower concentrations than the originally concentrated or pure chemicals. They should be informed of potential hazards and exercise cautions when dealing with chemicals. Appropriate personal protective equipment should be worn during the experiment. The alkaline oxygen− safranine−benzoin waste should be neutralized before disposal. Otherwise, small amounts of reaction wastes are safe to discard down the drain with copious amount of water.

■

RESULTS AND DISCUSSION Results from 14 groups of students from an introductory chemistry laboratory class during the third term of academic year 2017−18 are shown in Table 1. (Raw data and sample conclusions for the two reactions can be found in the Supporting Information.) • The first set of data in Table 2 is the order of reaction with respect to five reagents from the plot of natural logarithm of rate vs natural logarithm of concentration. This is comparable to the method of initial rate commonly found in general chemistry textbooks,45,46 but there are more data points in our experiments. The class mean values were close to the theoretical order of reaction value of 1 for benzoin, sodium hydroxide, potassium iodide and hydrogen peroxide, and 0 for cysteine. • The second-order rate constant k (in the unit of M−1 s−1) can be found from slopes and intercepts of graphs between rate vs concentration. Different approaches yield slightly different values of k, and students were encouraged to discuss possible reasons in their report. • Last, the plots of clock time vs concentration of cysteine or clock time vs 1/concentration of other reagents are linear. Most of students obtained r2 > 0.97. These r2 results provide confirmation of the pseudo-zeroth-order kinetic model used in the previous step.

■

CONCLUDING REMARKS We have designed two clock reaction experiments to support first-year undergraduate chemical kinetic teaching. The oxygen−safranin−benzoin clock is a variant of the popular blue bottle experiment, and the cysteine−iodine−hydrogen peroxide clock is a modification of the well-known iodine clock. Conducting the experiments in the same laboratory session allows students to clearly see the similarities and differences between the two reactions. The main point of this experiment is derivation of the rate law from an experimentally determined rate of reaction. Rates of reactions are derived from the initial concentration divided by the clock time. The initial rate method can be demonstrated to students without the need to measure concentrations of a reagent or a product using expensive equipment. The spreadsheet software helps students to obtain the numerical values of the order of reaction and the rate constant straightaway. Students therefore have more time to understand the derivation and the process to construct the spreadsheet file. Potential additional topics to be discussed during this experiment include a statistical test for regression coefficients, regression through the origin (RTO),52 nonlinear regression,53 mechanism of reaction, and redox reaction of biomolecules and catalysis.37 The statistical test and RTO have already been implemented in the worksheet. The clock reactions proposed here are also open to modifications; for example, it is well-known that benzoin/cysteine in the two clock reactions may be substituted by other appropriate reducing agents16 and the experiment may be scaled down to less than 1 mL on filter paper.42

One week after the experiment, students were given a hard copy of an anonymous one-page survey adapted from our previous version of a plus-minus-delta questionnaire51 (Supporting Information) with the graded laboratory report. The study was approved by the Institute for Population and Social Research’s Institutional Review Board, IPSR-IRB COA. No. 2016/12-150. Figure 2 shows tallies of the rating question collected from the students. These results indicate that students are generally satisfied with the two experiments (average score >4.0 for content, concept, and suitability). We acknowledge that kinetics is a rather difficult topic for most students; their average

■

ASSOCIATED CONTENT

S Supporting Information *

The Supporting Information is available on the ACS Publications website at DOI: 10.1021/acs.jchemed.8b00547. Instructor notes and laboratory manual (PDF, DOCX) Spreadsheet for calculations (XLSX, PDF)

Figure 2. Students’ rating (n = 30) of the four parts of the clock reaction experiments (1 = very unsatisfied and 5 = very satisfied). E



■

Communication

(14) Rajchakit, U.; Limpanuparb, T. Rapid Blue Bottle Experiment: Autoxidation of Benzoin Catalyzed by Redox Indicators. J. Chem. Educ. 2016, 93, 1490−1494. (15) Rajchakit, U.; Limpanuparb, T. Greening the Traffic Light: Air Oxidation of Vitamin C Catalyzed by Indicators. J. Chem. Educ. 2016, 93, 1486−1489. (16) Limpanuparb, T.; Roongruangsree, P.; Areekul, C. A DFT investigation of the blue bottle experiment: E◦ half‑cell analysis of autoxidation catalysed by redox indicators. R. Soc. Open Sci. 2017, 4 (11), 170708. (17) Chen, P. S. Entertaining and Educational Chemical Demonstrations; Chemical Elements Publishing Company: Camarillo, CA, 1974. (18) Chen, P. S. Resazurin-Reduction and Oxidation. J. Chem. Educ. 1970, 47, A335. (19) Cook, A. G.; Tolliver, R. M.; Williams, J. E. The Blue Bottle Experiment Revisited: How Blue? How Sweet? J. Chem. Educ. 1994, 71, 160−161. (20) Campbell, J. A. Why Do Chemical Reactions Occur?; Prentice-Hall: Englewood Cliffs, NJ, 1965. (21) Limpanuparb, T.; Areekul, C.; Montriwat, P.; Rajchakit, U. Blue Bottle Experiment: Learning Chemistry without Knowing the Chemicals. J. Chem. Educ. 2017, 94 (6), 730−737. (22) Snehalatha, T.; Rajanna, K.; Saiprakash, P. Methylene blueAscorbic acid: An undergraduate experiment in kinetics. J. Chem. Educ. 1997, 74, 228−233. (23) Church, J. A.; Dreskin, S. A. Kinetics of color development in the Landolt (″ iodine clock″) reaction. J. Phys. Chem. 1968, 72 (4), 1387− 1390. (24) Landolt, H. Ueber die Zeitdauer der Reaction zwischen Jodsäure und schwefliger Säure. Ber. Dtsch. Chem. Ges. 1886, 19 (1), 1317−1365. (25) McAlpine, R. The rate of oxidation of iodide ion by hydrogen peroxide. J. Chem. Educ. 1945, 22 (8), 387. (26) Sattsangi, P. D. A microscale approach to chemical kinetics in the general chemistry laboratory: The Potassium Iodide Hydrogen Peroxide Iodine-Clock reaction. J. Chem. Educ. 2011, 88 (2), 184−188. (27) Liebhafsky, H. A.; McGavock, W. C.; Reyes, R. J.; Roe, G. M.; Wu, L. S. Reactions involving hydrogen peroxide, iodine, and iodate ion. 6. Oxidation of iodine by hydrogen peroxide at 50. degree. C. J. Am. Chem. Soc. 1978, 100 (1), 87−91. (28) Liebhafsky, H. A.; Mohammad, A. The kinetics of the reduction, in acid solution, of hydrogen peroxide by iodide ion. J. Am. Chem. Soc. 1933, 55 (10), 3977−3986. (29) Sharma, K. R.; Noyes, R. M. Oscillations in chemical systems. 13. A detailed molecular mechanism for the Bray-Liebhafsky reaction of iodate and hydrogen peroxide. J. Am. Chem. Soc. 1976, 98 (15), 4345− 4361. (30) Seeger, W. Data for Time Reactions to Be Conducted by Students in the Laboratory. Journal of Chemical Education 1931, 8 (1), 166. (31) Wright, S. W. Tick tock, a vitamin C clock. J. Chem. Educ. 2002, 79 (1), 40A. (32) Wright, S. W. The vitamin C clock reaction. J. Chem. Educ. 2002, 79 (1), 41. (33) Wright, S. W.; Folger, M. R.; Rice, M. A. A Clock Reaction Sympathetic Ink from Consumer Chemicals. J. Chem. Educ. 2006, 83 (10), 1473. (34) Weinberg, R. B. An Iodine Fluorescence Quenching Clock Reaction. J. Chem. Educ. 2007, 84 (5), 797. (35) Vitz, E. A Student Laboratory Experiment Based on the Vitamin C Clock Reaction. J. Chem. Educ. 2007, 84 (7), 1156. (36) Pfennig, B. W.; Roberts, R. T. A Kinetics Demonstration Involving a Green-Red-Green Color Change Resulting from a LargeAmplitude pH Oscillation. J. Chem. Educ. 2006, 83 (12), 1804. (37) Copper, C. L.; Koubek, E. An Experiment to Demonstrate How a Catalyst Affects the Rate of a Reaction. J. Chem. Educ. 1999, 76 (12), 1714. (38) Carpenter, Y.-y.; Phillips, H. A.; Jakubinek, M. B. Clock Reaction: Outreach Attraction. J. Chem. Educ. 2010, 87 (9), 945−947.

Instructor presentation and videos of spectrophotometric measurement (ZIP) Survey instrument (PDF, DOCX)

AUTHOR INFORMATION

Corresponding Author

*E-mail: [email protected]. ORCID

Taweetham Limpanuparb: 0000-0002-8558-6199 Notes

The authors declare no competing financial interest.

■

ACKNOWLEDGMENTS We thank our full-time technician, Warangkana Yimkosol, and student assistants, Oranun Ongwisespaiboon, Cherprang Areekul, Yonlada Nawilaijaroen, and Rattha Noorat, for their continuous and diligent support. We appreciate the invaluable feedback and comments from students in the ICCH224 class during and after the experiments. Generous equipment and material supplies from Mahidol University International College (MUIC) and the Institute for the Promotion of Teaching Science and Technology (IPST) are also acknowledged.

■

REFERENCES

(1) Schreiner, R.; Shakhashiri, B. Z.; Scott, E. S.; Bell, J. A.; Testen, M. E. Clock reaction. In Chemical Demonstrations; Shakhashiri, B. Z., Ed.; The University of Wisconsin Press: Madison, WI, 1992; Vol. 4, pp 1− 89. (2) Lente, G.; Bazsa, G.; Fábián, I. What is and what isn’ta clock reaction? New J. Chem. 2007, 31 (10), 1707−1707. (3) Horváth, A. K.; Nagypál, I. Classification of clock reactions. ChemPhysChem 2015, 16 (3), 588−594. (4) Xu, L.; Horváth, A. K. A Possible Candidate to Be Classified as an Autocatalysis-Driven Clock Reaction: Kinetics of the Pentathionate− Iodate Reaction. J. Phys. Chem. A 2014, 118 (32), 6171−6180. (5) Xu, L.; Horváth, A. K. Autocatalysis-Driven Clock Reaction II: Kinetics of the Pentathionate−Periodate Reaction. J. Phys. Chem. A 2014, 118 (42), 9811−9819. (6) Cseko, G.; Valkai, L.; Horváth, A. K. A Simple Kinetic Model for Description of the Iodate−Arsenous Acid Reaction: Experimental Evidence of the Direct Reaction. J. Phys. Chem. A 2015, 119 (45), 11053−11058. (7) Valkai, L.; Cseko, G.; Horváth, A. K. Initial inhomogeneityinduced crazy-clock behavior in the iodate−arsenous acid reaction in a buffered medium under stirred batch conditions. Phys. Chem. Chem. Phys. 2015, 17 (34), 22187−22194. (8) Valkai, L.; Horváth, A. K. Compatible Mechanism for a Simultaneous Description of the Roebuck, Dushman, and Iodate− Arsenous Acid Reactions in an Acidic Medium. Inorg. Chem. 2016, 55 (4), 1595−1603. (9) Campbell, J. KineticsEarly and often. J. Chem. Educ. 1963, 40, 578−583. (10) Dutton, F. Methylene blue-Reduction and oxidation. J. Chem. Educ. 1960, 37, A799. (11) Anderson, L.; Wittkopp, S. M.; Painter, C. J.; Liegel, J. J.; Schreiner, R.; Bell, J. A.; Shakhashiri, B. Z. What Is Happening When the Blue Bottle Bleaches: An Investigation of the Methylene BlueCatalyzed Air Oxidation of Glucose. J. Chem. Educ. 2012, 89, 1425− 1431. (12) Engerer, S. C.; Cook, A. G. The blue bottle reaction as a general chemistry experiment on reaction mechanisms. J. Chem. Educ. 1999, 76, 1519−1520. (13) Chen, P. S. Oxidation and Reduction of Indigo Carmine Aided by Benzoin. J. Chem. Educ. 1970, 47, A335. F



Communication

(39) Burgess, A. E.; Davidson, J. C. Kinetics of the Rapid Reaction between Iodine and Ascorbic Acid in Aqueous Solution Using UV− Visible Absorbance and Titration by an Iodine Clock. J. Chem. Educ. 2014, 91 (2), 300−304. (40) Bauer, J.; Tomisic, V.; Vrkljan, P. B. A. Kinetics and Mechanism of Iodide Oxidation by Iron(III): A Clock Reaction Approach. J. Chem. Educ. 2008, 85 (8), 1123. (41) Bauer, J.; Tomišić, V.; Vrkljan, P. B. A. The Effect of Temperature and Ionic Strength on the Oxidation of Iodide by Iron(III): A Clock Reaction Kinetic Study. J. Chem. Educ. 2012, 89 (4), 540−544. (42) Barrera, L. A.; Escobosa, A. C.; Alsaihati, L. S.; Noveron, J. C. Conducting a Low-Waste Iodine Clock Experiment on Filter Paper To Discern the Rate Law. J. Chem. Educ. 2019, 96 (1), 165−168. (43) Henary, M. M.; Russell, A. A. An Inexpensive Kinetic Study: The Reaction of FD&C Red #3 (Erythrosin B) with Hypochlorite. J. Chem. Educ. 2007, 84 (3), 480. (44) Ealy, J. B.; Negron, A. R.; Stephens, J.; Stauffer, R.; Furrow, S. D. The Glyoxal Clock Reaction. J. Chem. Educ. 2007, 84 (12), 1965. (45) Petrucci, R. H.; Herring, F. G.; Bissonnette, C.; Madura, J. General Chemistry: Principles and Modern Applications; Pearson Canada Inc.: Ontario, Canada, 2016. (46) Zumdahl, S. S.; Zumdahl, S. A.; DeCoste, D. J. Chemistry; Cengage Learning: Boston, MA, 2016. (47) Lim, K. F. Negative pH Does Exist. J. Chem. Educ. 2006, 83 (10), 1465. (48) Flinn Scientific. Reaction Kinetics in Blue. https://www.youtube. com/watch?v=InhTcvg1AlA (accessed Feb 2019). (49) Rushton, G. T.; Criswell, B. A.; McAllister, N. D.; Polizzi, S. J.; Moore, L. A.; Pierre, M. S. Charting an Alternate Pathway to Reaction Orders and Rate Laws in Introductory Chemistry Courses. J. Chem. Educ. 2014, 91 (1), 66−73. (50) Casado, J.; Lopez-Quintela, M. A.; Lorenzo-Barral, F. M. The initial rate method in chemical kinetics: Evaluation and experimental illustration. J. Chem. Educ. 1986, 63 (5), 450. (51) Limpanuparb, T.; Kanithasevi, S.; Lojanarungsiri, M.; Pakwilaikiat, P. Teaching Boyle’s Law and Charles’ Law through Experiments that Use Novel, Inexpensive Equipment Yielding Accurate Results. J. Chem. Educ. 2019, 96 (1), 169−174. (52) Eisenhauer, J. G. Regression through the origin. Teaching statistics 2003, 25 (3), 76−80. (53) Perrin, C. L. Linear or Nonlinear Least-Squares Analysis of Kinetic Data? J. Chem. Educ. 2017, 94 (6), 669−672.

G


Clock Reaction Revisited: Catalyzed Redox Substrate-Depletive

Recommend Documents