Laboratory Experiment pubs.acs.org/jchemeduc
A Stopped-Flow Kinetics Experiment for the Physical Chemistry Laboratory Using Noncorrosive Reagents Richard V. Prigodich* Chemistry Department, Trinity College, Hartford, Connecticut 06106-3100, United States S Supporting Information *
ABSTRACT: Stopped-flow kinetics techniques are important to the study of rapid chemical and biochemical reactions. Incorporation of a stopped-flow kinetics experiment into the physical chemistry laboratory curriculum would therefore be an instructive addition. However, the usual reactions studied in such exercises employ a corrosive reagent that can over time cause damage to sensitive and expensive components of a stopped-flow apparatus. The reduction of the dye 2,6-dichlorophenolindophenol by L-ascorbic acid at a pH of 8.5 is a reaction that provides accurate and reproducible data in the physical chemistry laboratory. Furthermore, these reagents are noncorrosive and protect instrumentation from damage.
KEYWORDS: Upper-Division Undergraduate, Physical Chemistry, Laboratory Instruction, Kinetics
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Table 1. Acidity Constants of the Reactants and Products of the Reaction between L-Ascorbic Acid and 2,6Dichlorophenolindophenol5
inetics experiments are crucial to complement the lecture coverage of this core topic in junior/senior year physical chemistry courses. There are many very good experiments that can be employed in this portion of the laboratory syllabus to compare different mathematical methods to analyze kinetics data, or examine the primary kinetic salt effect, for example.1,2 Stopped-flow kinetics studies are sometimes included in physical chemistry laboratory courses, if such instrumentation is available. This is an excellent way to introduce students to an advanced physical chemical technique and to demonstrate the fact that chemical reaction rates cover many orders of magnitude. One difficulty is that the classic stopped-flow experiments often utilize corrosive reagents, such as perchloric acid.3,4 The recent acquisition of an instrument that was capable of circular dichroism spectroscopy as well as stopped-flow experiments made inclusion of a stopped-flow experiment in our physical chemistry laboratory curriculum feasible. An inquiry to the manufacturer of the instrument regarding suggestions for suitable reactions for stopped-flow experiments using noncorrosive reagents produced a reference to a paper that is much cited for the determination of the dead time of a stopped-flow apparatus.5 One of the reactions suggested for this calibration is the reduction of 2,6-dichlorophenolindophenol by L-ascorbic acid, two noncorrosive water-soluble reagents. For example, in 1976 this reaction was used to demonstrate the capabilities of a home-built spring-loaded stopped-flow apparatus.6 At first glance the reaction may seem problematical for an undergraduate experiment due to the multiple ionization states possible for the reactants (Table 1). © XXXX American Chemical Society and Division of Chemical Education, Inc.
Reactant L-Ascorbic
acid 2,6-Dichlorophenolindophenol 2,6-Dichlorophenolindophenol (reduced form)
Ka1
Ka2
9.16 × 10−5 3.0 × 10−1 1 × 10−7
4.57 × 10−12 2.69 × 10−6 none
Thus, at many pH’s two or more reactions may be occurring simultaneously. However, at pH = 8.5, only single ionization states exist for each chemical species. The data from kinetics studies of this reaction are consistent with a two-step reaction mechanism.5 The first and rate-determining step is the reaction of ascorbic acid with 2,6-dichlorophenolindophenol to form radical species of each reactant. In the second step the radicals rapidly react to form products. Protons can be produced in each step. The numbers depend on pH and hence the requirement for a buffer. At pH = 8.5 the second-order rate constant is 2.99 × 102 M−1 s−1. Therefore, the reaction at a pH of 8.5 would allow for straightforward analysis of the data and is appropriate for an undergraduate physical chemistry laboratory stopped-flow experiment. The reaction should be done under pseudo-firstorder conditions at high L-ascorbic acid concentrations. Since students are taught the method of initial rates to determine the reaction orders of reactants in chemical reactions in introductory chemistry courses, application of the method of initial rates to this reaction is pedagogically useful. Analysis of
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initial rate data (over the first 50 ms) with various combinations of reactant concentrations will allow students to determine the rate order for each reactant. Once the orders of reaction are determined, appropriate (first-order) analysis of a full onesecond data set will allow for the determination of the pseudofirst-order rate constant kobs and the true second-order rate constant (k2), where kobs = k2[L-ascorbic acid].
HAZARDS All reagents are water-soluble and pose no special hazard. The phosphate buffer and dye are irritants, however, and nitrile gloves are to be worn by students. The reaction solutions and excess reagents containing 2,6-dichlorophenolindophenol are pooled as chlorinated hydrocarbons and excess buffers, and ascorbic acid solutions are pooled as aqueous waste. The chemical hygiene officer disposes of these solutions in the appropriate manner.
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EXPERIMENTAL DETAILS The recommended reaction solutions are given in the Supporting Information. It is also recommended that L-ascorbic acid sodium salt (ACROS) be purchased new for these reactions. The sodium salt of 2,6-dichlorophenolindophenol (ACROS) was used and stored in a desiccator. All other chemicals are reagent grade. All solutions were made with distilled deionized water. A stopped-flow instrument capable of recording absorbance data at 524 nm over a 50 ms time scale is needed to determine initial rates. The ability to perform the reactions at a constant selected temperature (25 °C) will allow for comparison to literature values of the second-order rate constant, but is not necessary. Alternatively, less expensive apparatus are available that have longer dead times. Using such systems analysis of absorbance versus time data on a seconds time scale is a viable option. By plotting the natural log of absorbance versus time and the inverse of absorbance versus time students will be able to determine the reaction orders and then calculate the true second-order rate constant. It is convenient if the software that operates the stopped-flow apparatus has fitting routines for linear data sets to determine initial rates and for exponential decays to determine the observed first-order rate constant under pseudo-first-order conditions. However, again, this is not strictly necessary, and data analysis can be performed by other means. The instrument used was an OLIS conversion of a Cary 16 spectrophotometer to a stopped-flow apparatus. The instrument parameters were the following: the monochromator was set to 524 nm, the slit width was set at 1 mm, and the collection time was set for 1.0 s. 0.25 mL samples of each reactant were injected into the stopped-flow cell. The instrument design was such that the cell path length was 20 mm. The cell volume was 35 μL. The drive syringes were kept at 25.0 °C with a Fisher IsoTemp recirculating bath attached to the sample chamber. The first 50 ms of the data in each run was used to determine the initial rate of reaction by fitting this data to a linear equation, the slope being the initial rate. The reaction order of each reactant is determined by comparing initial rates from a set of reactions in which the concentration of one reactant is held constant while the other reactant concentration varies according to eq 1. [R] is the concentration of the reactant being varied, υi is the initial rate, and the subscripts 1 and 2 refer to the initial rates of reaction at specific and different [R]. reaction order =
log([R]1 /[R]2 ) log(υi,1/υi,2)
Laboratory Experiment
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RESULTS AND DISCUSSION The pairs of solutions reacted and the kinetics parameters derived from these reactions are given in Table 2. The secondTable 2. Kinetics Parameters Determined for the Reaction of 2,6-Dichlorophenolindophenol (DCPIP) and L-Ascorbic Acid (L-AA)a [L-AA] (M)
[DCPIP] (M)
k2 (M−1 s−1)
kobs (s−1)
0.0333 0.000165 5.15 ± 0.06 310. ± 4 [L-AA] (M) [DCPIP] (M) Initial Rate (s−1) 0.0333 0.0333 0.0333 0.00888 0.00444 0.00222 a
0.000165 0.000110 0.000055 0.000165 0.000165 0.000165
3.05 2.02 1.11 9.71 5.17 2.53
± ± ± ± ± ±
0.06 0.01 0.12 0.26 0.30 0.01
Literature Value for k24 (M−1 s−1) 299 Reaction Order DCPIP 1.08 ± 0.10
L-AA
1.03 ± 0.05
Note: The kinetics parameters are the averages of three runs.
order rate constant determined in these studies is close to the literature value. Likewise, the reaction orders for the two reactants are within experimental error of the expected firstorder result. The reaction solutions are not hazardous or corrosive. The stopped-flow apparatus can be easily cleaned with distilled deionized water. The data obtained in these reactions is reproducible and accurate in the hands of junior/senior level chemistry students. This laboratory exercise can be easily completed in the usual 3 h laboratory period, especially if the 10× stock buffer and 1× buffer are prepared beforehand. The actual reaction solutions should be made during the laboratory period; however, it is essential that ascorbic acid solutions be prepared immediately before each kinetics run. The experiment is probably best done in groups of 2−3 students. This experiment provides an effective means of introducing students to the utility of stopped-flow experiments as well as providing an exercise in determining second-order rate constants and reactant reaction orders. At the start of the experiment students can mix 1 mL each of the reactants to observe how apparently instantaneously the reaction occurs. An optional additional exercise would be to have students take spectra of the oxidized and reduced forms of 2,6-dichlorophenolindophenol to determine an optimal wavelength for analysis.
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(1)
The full one-second data set was fit to a single exponential decay routine that yields the kobs for the pseudo-first-order reactions. The second-order rate constant is calculated by dividing kobs by the constant L-ascorbic acid concentration (half the concentration injected).
ASSOCIATED CONTENT
S Supporting Information *
A complete description of the laboratory including instructions for students performing the lab. This material is available via the Internet at http://pubs.acs.org. B
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Laboratory Experiment
AUTHOR INFORMATION
Corresponding Author
*E-mail:
[email protected]. Notes
The authors declare no competing financial interest.
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ACKNOWLEDGMENTS The author wishes to thank the staff at Olis, Inc., for the suggestion of this reaction for this laboratory exercise and their helpful technical advice.
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REFERENCES
(1) Hemelatha, M. R. K.; Noor-Botcha, I. An Undergraduate Physical Chemistry Experiment on the Analysis of First-Order Kinetic Data. J. Chem. Educ. 1997, 74 (8), 972−974. (2) Watkins, K. W.; Olson, J. A. Ionic Strength Effect on the Rate of Reduction of Hexacyanoferrate(III) by Ascorbic Acid. J. Chem. Educ. 1980, 57 (2), 158−159. (3) Clark, C. R. A Stopped-Flow Kinetics Experiment for Advanced Undergraduate Laboratories: Formation of Iron(III) Thiocyanate. J. Chem. Educ. 1997, 74 (10), 1214−1217. (4) Hoag, C. M. Simple and Inexpensive Computer Interface to a Durrum Stopped-Flow Apparatus Tested Using the Iron(III)− Thiocyanate Reaction. J. Chem. Educ. 2005, 82 (12), 1823−1825. (5) Tonomura, B.; Kakatani, H.; Ohnishi, M.; Yamaguci-Ito, J.; Hiromi, K. Test Reactions for a Stopped-Flow Apparatus. Anal. Biochem. 1978, 84, 370−383. (6) Morelli, B. A Kinetic Experiment Using a Spring Powered, Stopped-Flow Apparatus. J. Chem. Educ. 1976, 53 (2), 119−122.
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