Reaction of FD&C Blue 1 with Sodium Percarbonate: Multiple Kinetics

Laboratory Experiment Cite This: J. Chem. Educ. 2019, 96, 1453−1457

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Reaction of FD&C Blue 1 with Sodium Percarbonate: Multiple Kinetics Methods Using an Inexpensive Light Meter Ruth E. Nalliah* Department of Chemistry, Huntington University, 2303 College Avenue, Huntington, Indiana 46750, United States

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S Supporting Information *

ABSTRACT: In selecting from a repertoire of traditional kinetics experiments, an instructor often has to choose among having students gain experience with the graphical method, the method of initial rates, or a temperature-dependent experiment in which students construct an Arrhenius plot. This paper presents an environmentally friendly bleaching reaction between the dye FD&C Blue 1 and a nonchlorinated bleach alternative, sodium percarbonate, which goes to completion within 15 min and can be monitored with an inexpensive light meter. The fast time scale of the reaction allows students to use the graphical method along with the method of initial rates to determine the overall rate law, as well as collect data for an Arrhenius plot, potentially within a 3 h laboratory period. Students can generate an Arrhenius plot quickly and inexpensively by detecting the extent of reaction in a beaker using a desk lamp, filter, and light meter, eliminating the need for spectrophotometer access, a jacketed cell, and a heated circulating bath. Performing all three types of kinetics experiments with one simple reaction in one or two laboratory periods allows students to see how the methods are interrelated and complementary. KEYWORDS: First-Year Undergraduate/General, High School/Introductory Chemistry, Laboratory Instruction, Hands-On Learning/Manipulatives, Consumer Chemistry, Dyes/Pigments, Laboratory Equipment/Apparatus, Oxidation/Reduction, Rate Law, UV−Vis Spectroscopy

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INTRODUCTION Previous student kinetics experiments have been developed which involve the degradation of food dyes with hypochlorite,1,2 while this paper describes a kinetics experiment using a nonchlorinated bleaching agent, sodium percarbonate, Na2CO3·1.5 H2O2, to degrade the food dye FD&C Blue 1 (Blue 1). The mixing of hydrogen peroxide alone with Blue 1 does not result in significant bleaching within an appreciable time scale; however, sodium percarbonate simultaneously creates basic solution conditions and releases hydrogen peroxide, which bleaches the dye as reported for similar chromophores.3−5 The experiment is of value in discussing nonchlorinated bleach alternatives, as well as their usage in the oxidative degradation of pollutants such as more toxic dyes and pharmaceuticals.6 The bleaching reaction with sodium percarbonate gives a definite first-order kinetics plot with respect to Blue 1 and goes to completion in about 15 min under the conditions that we use here. The reaction is quick, convenient, environmentally friendly, inexpensive, visually appealing, and ideal for maximizing the use of laboratory time. The reaction can be used for students to experience three major types of kinetics experiments in one or two laboratory periods: the graphical method, the method of initial rates, and the construction of an Arrhenius plot. The value for students using all three types of methods to investigate one reaction is that they are more apt to observe how these methods relate to © 2019 American Chemical Society and Division of Chemical Education, Inc.

each other in characterizing the rate of a reaction. An integrated kinetics experiment has also been developed for the dissociation of ferroin,7 but with more extensive reliance on a spectrophotometer. Inexpensive colorimeter circuits as well as smartphone setups have been described previously as economical ways to measure solution concentrations without commercial spectrophotometers.8−11 The experiment described in this paper involves measuring the degradation of the dye using inexpensive light meters with 50 mL solution quantities, which offer the advantage of greater simplicity as well as eliminate the need for water baths, circulators, and jacketed cells for temperature-dependent measurements. Other creative methods which have been used for student measurement of activation energies as documented in this Journal range from gravity-powered flow systems, to magnetic levitation, to water displacement.12−19 The development of efficient, simple, and inexpensive methods for students to measure activation energies continues to be a challenge, and instructors may find our method useful. The rate of this reaction is expressed as Received: July 23, 2018 Revised: May 30, 2019 Published: June 19, 2019 1453

DOI: 10.1021/acs.jchemed.8b00589 J. Chem. Educ. 2019, 96, 1453−1457

Journal of Chemical Education rate = k[Blue 1]x [H 2O2 ] y

Laboratory Experiment

cleaning agents such as powdered stain removers, calculating the exact concentrations of the reactants becomes a challenge with commercial substances. It should be noted that a pocket light meter which is manufactured with a white translucent dome over the detector, which helps diffuse the incoming light, gives more reproducible results for this experiment than a light meter without a diffuser dome, or an app using a smartphone without a diffuser. In addition, the apparatus could not be used to calculate the correct absorbances unless the filter were to pass a more narrow wavelength range than the absorption band of the solution; otherwise, calculated absorbances will be too low. Regardless, the apparatus can be calibrated with a reference solution to determine the transmitted light intensity corresponding to a given percentage of dye degradation, and it also allows students to observe that greater amounts of red light pass through less concentrated blue solutions.

(1)

in which k is the rate constant, concentrations of each species are indicated by brackets, x is the order of reaction with respect to Blue 1, and y is the order of reaction with respect to the H2O2 released from sodium percarbonate. Previous kinetics literature in this Journal stresses the importance of students learning several different experimental methods for determining the orders of reaction, including the advantages and drawbacks of each method. Students should realize that the method of initial rates is useful for determining the order of previously uncharacterized reactions, due to possible side reactions and other effects that can become more pronounced as the reaction proceeds.20,21 The graphical method of plotting concentration data as a function of time using the integrated rate equations can be used to confirm the order of reaction and to get a reliable determination of the rate constant.20 In order to characterize the reaction between Blue 1 and hydrogen peroxide, we use the graphical method to determine the reaction order with respect to the Blue 1 dye (expressed as x in eq 1), and the method of initial rates to determine the reaction order with respect to the hydrogen peroxide released from the sodium percarbonate (expressed as y in eq 1). With the graphical determination of the reaction order with respect to Blue 1, concentrations are such that the hydrogen peroxide produced from the sodium percarbonate is over 1000 times in excess of the Blue 1 concentration, so that the hydrogen peroxide concentration can be considered approximately constant over the course of the reaction. When the graphical method is used, collecting data for 4−5 half-lives is recommended in order to be able to reliably determine the order of reaction.20 For the reaction described here, collection of data for approximately 5 half-lives can be completed within 15 min. In this experiment, the order of reaction with respect to hydrogen peroxide is measured by approximating the method of initial rates. For the design of student kinetics experiments, the method of initial rates should be designed to measure the initial rate of the reaction, rather than the time required for completion.21 With more precise measurement methods, the method of initial rates may involve calculating the initial slope of concentration data obtained during the first 10% of the reaction.21 In this reaction, we approximate the method of initial rates by measuring the time required for the reaction to go 20% to completion, in order for students to obtain a more measurable difference in the intensity of red light transmitted by the dye in solution as measured by an inexpensive light meter. Finally, the method of initial rates is used to measure the initial rates of the reaction at different temperatures, in order to estimate the activation energy of the reaction using the Arrhenius equation.

Part 1: Graphical Method: Analysis of the Reaction Order with Respect to Blue 1

If spectrophotometers are available, students mix a small amount of sodium percarbonate solution with Blue 1 solution directly in the cuvette and obtain the absorbance at approximately 630 nm as a function of time for up to 15 min, or until the absorbance goes to zero. The data can be collected automatically in kinetics mode with digital spectrophotometers. If spectrophotometers are not available, students can use the light meters (Figure 1) to measure the

Figure 1. Diagram of the apparatus used to measure the red light transmitted through the solution. The solution is tested in the beaker, which can also be insulated with bubble wrap if needed for temperature-dependent experiments.

light intensities transmitted through calibration solutions (such as 100%, 80%, 60%, 40%, 20%, 10%, and 5% of the original concentration of the Blue 1 stock solution) and monitor the time it takes for the transmitted light intensity to increase to each of those benchmark levels during the reaction. Part 2: Method of Initial Rates: Analysis of the Reaction Order with Respect to Hydrogen Peroxide

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EXPERIMENTAL OVERVIEW Items needed for this experiment beyond routine labware include light meters (under $20 each, ideally one per student pair), commercial desk lamps with appropriate bulbs, red plastic filters, and either hot plates or laboratory-grade microwave ovens. Digital spectrophotometers are a plus for the graphical method but are not required. Chemicals needed include FD&C Blue 1 dye and sodium percarbonate. Although the reaction can also be done with commercial food coloring solutions and commercial sodium percarbonate containing

After the concept of absorbance is discussed, students set up the light meter apparatus as shown in Figure 1. To obtain a reference intensity for determining when the reaction has gone 20% to completion, the students dilute a Blue 1 stock solution to 80% of its original concentration and measure the light intensity transmitted through 50 mL of the solution in a beaker. To begin the reaction, students add sodium percarbonate solution to 50 mL of undiluted Blue 1 stock solution in an identical beaker and determine how long it takes for 20% of the Blue 1 dye to react, as measured by how long it 1454


Journal of Chemical Education


Figure 2. Zero-order, first-order, and second-order plots for the reaction between Blue 1 and sodium percarbonate which show first-order kinetics with respect to the Blue 1 dye. The first row of plots shows student data collected by measuring the absorbance of Blue 1 at 628.5 nm with a spectrophotometer in kinetics mode; the second row of plots shows student data obtained using a light meter. The concentration of sodium percarbonate was approximately 0.024 M, giving a hydrogen peroxide concentration of approximately 0.036 M.

solutions are mixed with heated solutions. Students should make and keep their sodium percarbonate stock solutions in a fume hood. All mixing and analysis of heated solutions need to be done in a fume hood, including analysis with the light meter. For this particular reaction, this precaution may limit the scope of the experiment for laboratories not equipped with fume hoods, or for online laboratories done in home settings.

takes for the transmitted light intensity to increase to the level observed for the diluted stock solution. The procedure is repeated with varying initial concentrations of sodium percarbonate to determine the order of reaction with respect to the hydrogen peroxide released from the sodium percarbonate. Part 3: Data Collection for an Arrhenius Plot

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Students measure initial rates as described in Part 2 at room temperature and three additional temperatures, under identical concentration conditions, to obtain a total of four data points for an Arrhenius plot. Placing a beaker of 50 mL of dye solution in an ice bath or heating it on a hot plate for several minutes quickly generates temperature ranges from about 15 to 32 °C prior to addition of sodium percarbonate solution. If desired, heat gain or loss can be minimized by using bubble wrap to insulate the cylindrical part of the beaker, although during the time scale of the initial rate measurement for this reaction, heat gain or loss is minimal, even in uninsulated beakers. At temperatures higher than 45 °C, the reaction becomes too fast for students to monitor reliably with respect to the initial rate.

RESULTS AND DISCUSSION Figure 2 compares concentration data in zero-, first-, and second-order plots used to determine the order of the reaction with respect to Blue 1, as measured with spectrophotometers and with light meters. In each case, the linearity of the plot of the natural log of Blue 1 concentration vs time gives a clear indication of a first-order reaction with respect to Blue 1 (the average R2 value for a first-order plot was 0.997 with a standard deviation of 0.003 for 7 student pairs using spectrophotometers, and 0.988 with a standard deviation of 0.019 for 8 student pairs using light meters). Table 1 shows data for the method of initial rates used to determine the reaction order with respect to the hydrogen peroxide released from the sodium percarbonate. When the initial concentration of Blue 1 is held constant, the rate is directly proportional to the initial concentration of hydrogen peroxide, indicating that the reaction is first-order with respect to hydrogen peroxide; data collected by 19 pairs of students in our laboratory has generated an average order of 1.1 with a standard deviation of 0.3. A first-order reaction both with respect to the dye and with respect to the hydrogen peroxide released from the sodium percarbonate is consistent with that

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HAZARDS AND PRECAUTIONS Students should be warned to notify the instructor of any food coloring allergies. The sodium percarbonate may cause irritation. More importantly, the hydrogen peroxide vapors given off from concentrated sodium percarbonate solutions may cause mild respiratory irritation in the days following the experiment, which may not be noticed during the experiment itself. This is especially pronounced when sodium percarbonate 1455




first week, and Parts 2 and 3 in the second week, with some laboratory time devoted to students processing their data, checking the sensibility of their calculations, and commenting on the outcomes. The method of initial rates, especially when performed at temperatures other than room temperature, has several time-sensitive procedures which students must choreograph and rehearse, because they cannot stop to think about what to do next. The process can be an exciting challenge, especially for premedical students thinking about scenarios such as critical surgeries, which must also be rehearsed. At our institution, the format of this experiment had been under development for a number of years before the light meters were found to be an ideal tool to measure the color fading for the bleaching reaction at different temperatures. In a second semester principles of chemistry course for science majors, students performed the experiment during a single 3 h laboratory period for two years, with 9 and 14 students, respectively, working in pairs, after having done similar kinetics problems in the lecture portion of the course. In the third year, a group of 15 students completed Part 1 with light meters in a 3 h laboratory period in week 1, and Parts 2 and 3 in a 3 h laboratory period in week 2. With this format, they seemed to be less overwhelmed when processing the data, and lab report scores improved (see Supporting Information). The student laboratory reports included an introduction; a zero-order, first-order, and second-order plot; the method of initial rates calculations; and an Arrhenius plot, as well as answers to experiment questions and conclusions. As a result of the experiment, students were able to compare practical advantages and disadvantages of the method of initial rates vs the graphical method of analysis. Most students indicated that they preferred the graphical method because, in this experiment, they felt that it was less susceptible to multiple sources of experimental error. While becoming aware of the importance of experimental consistency in using the method of initial rates, students can typically be proud of the fact that their initial rate data indicates a reaction order that is close to first-order, and that their reaction rates clearly speed up with increasing temperature, allowing them to calculate an activation energy. After monitoring the color disappearance under carefully controlled conditions in this experiment, students were able to make connections between the rate law determination methods used in most textbooks and their practical importance in addressing research problems such as determining the most efficient conditions for degrading organic pollutants.

Table 1. Student-Generated Data from the Method of Initial Rates Using Light Meters, 20.9 °C Molarity of H2O2a

Molarity of Blue 1

Time for 20% of Blue 1 to React (s)

0.0425 0.0212 0.0106

1.88 × 10−5 1.88 × 10−5 1.88 × 10−5

78 146 308

Initial Rate (M/s)

Calculated Order with Respect to H2O2b

4.82 × 10−8 2.57 × 10−8 1.22 × 10−8

0.904 0.990

a

The molarity of hydrogen peroxide in solution was calculated from that of the sodium percarbonate. bEach order shown here was calculated by using the data for two different H2O2 concentrations. The order can also be calculated by plotting ln(initial rate) vs ln([H2O2]) and taking the slope.

reported in the literature for the bleaching of other dyes with hydrogen peroxide under alkaline conditions.3,4 Figure 3 and Table 2 show student Arrhenius data for this reaction. The data are nearly consistent with the general rule of

Figure 3. Student-generated Arrhenius plot for the reaction between Blue 1 and sodium percarbonate, obtained by measuring the initial rates using a light meter at temperatures between 16.6 and 30.9 °C.

Table 2. Student Arrhenius Data from the Method of Initial Rates Using Light Meters Temperature (°C)

Estimated k (M−1 s−1)

16.6 21.9 25.0 30.9

0.0533 0.0802 0.111 0.171

thumb that many reaction rates double with a 10 °C increase in temperature. The student data shown in Figure 3 and Table 2 give an estimated activation energy of 61 kJ/mol; student data sets collected thus far with an R2 value greater than 0.81 have given an average activation energy of 59 with a standard deviation of 15 kJ/mol (n = 16, average R2 = 0.94), which is within range of the activation energies reported for the oxidation of phenolphthalein, alizarin, and crocetin by hydrogen peroxide under basic conditions (43.3, 68.8, and 69.9 kJ/mol, respectively).3,4 If laboratory-grade microwave ovens are available, we find that students can finish all parts of the experiment within a 3 h laboratory period which includes 20 min of briefing at the beginning; however, this procedure leaves very little time for students to begin processing their data during the laboratory period. The experiment leaves students with an abundance of kinetic data which may be overwhelming. For this reason, it may be most beneficial to have students complete Part 1 in the

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CONCLUSIONS

A kinetics experiment has been developed using a quick reaction, simple and inexpensive equipment, minimal solution preparation, and minimal hazards, in which students use the method of initial rates and the graphical method to determine the complete rate law and activation energy of a reaction within one or two laboratory periods. The experiment tends to be well-received by students and allows them to make connections to environmental applications of kinetic measurements. The light meter method of monitoring a reaction rate at various temperatures can be extended to other reactions containing a chromophore. 1456



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ASSOCIATED CONTENT

S Supporting Information *

The Supporting Information is available on the ACS Publications website at DOI: 10.1021/acs.jchemed.8b00589. Instructor information including purchasing information, preparation and briefing recommendations, answers to questions, experiment extensions, and other experiment information (PDF, DOCX) Student assessment (PDF, DOCX) Student instructions for laboratories with spectrophotometers (PDF, DOCX) Student instructions for laboratories with no spectrophotometers (PDF, DOCX)

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AUTHOR INFORMATION

Corresponding Author

*E-mail: [email protected]. ORCID

Ruth E. Nalliah: 0000-0002-2420-1734 Notes

The author declares no competing financial interest.

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ACKNOWLEDGMENTS This work was supported by the Emeriti Fellows Research and Artistic Creation Fund of Huntington University. The author would like to thank the undergraduate chemistry students at Huntington University for their eager participation in the development of this experiment.

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REFERENCES

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Reaction of FD&C Blue 1 with Sodium Percarbonate: Multiple Kinetics

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