In the Laboratory
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Decomposition of Aspartame A Kinetics Experiment for Upper-Level Chemistry Laboratories Kathryn R. Williams,* Bhavin Adhyaru, Julia Timofeev, and Michael Keith Blankenship Department of Chemistry, University of Florida, Gainesville, FL 32611-7200; *
[email protected] The artificial sweetener α-L-aspartyl-L-phenylalanine-1methyl ester, commonly called aspartame (Figure 1A), has been the subject of a number of undergraduate experiments, including several analyses (1–4), a synthesis (5), and a sweetness response study (6). This article describes a kinetics study of aspartame decomposition in aqueous solution. The experiment is suitable for the physical and biophysical chemistry laboratories or for use as a followup study to one of the analyses referenced above. Considering the widespread use of this compound in diet beverages, yogurt, and ice cream, there is high student interest in the decomposition process. The article by Stein (6) includes a summary of information about possible side effects of aspartame. A more complete review of the literature is given by Stamp (7). Description of the Experiment Aspartame decomposes in aqueous solution to a variety of products depending on the pH (8–11). In the neutral and basic region, the predominant route is loss of methanol to form 3-carboxymethyl-6-benzyl-2,5-dioxopiperazine (Figure 1B), usually shortened to 2,5-diketopiperazine or DKP, and some α-L-aspartyl-L-phenylalanine. In the student experiment, the pH is maintained at 7.0, where the primary product is DKP, and aspartame exists as an approximately 50:50 mixture of the zwitterion and anion forms. The presence of the terminal amino group in the deprotonated form is necessary for nucleophilic attack on the carboxyl carbon, as shown in Scheme I (11). In addition to the pH dependence, the rate is also affected by the identity and concentration of the buffer system. In particular, there is a major rate enhancement when phosphate is present, presumably owing to its capability to donate and accept a proton simultaneously, as shown in Scheme I (11). In this experiment, the effect of buffer identity is investigated by measuring the rates in 0.20 M phosphate and 0.20 M citrate, both of which are common ingredients in diet beverages, but
at much lower concentrations in the commercial products. The phosphate also contains NaCl to raise the total ionic strength to 1.0 M, which is the same as that of 0.20 M citrate. Students also determine the rates at two temperatures, 40 ⬚C and 55 ⬚C, to evaluate the activation energy, which they use to calculate the rate constant at room temperature. These parameters (neutral pH, high buffer concentration, and elevated temperature) were chosen to allow completion of the experiment within a five-hour laboratory period. The three systems—0.20 M citrate at 40⬚, 0.20 M phosphate at 40⬚, and 0.20 M phosphate at 55⬚—have half-lives of approximately 248 min, 38 min, and 11 min, respectively. In contrast, the half-life is 49 hours at pH 7.0 and 25 ⬚C with a phosphate concentration of 0.01M (11), which is still roughly three times the concentration in a typical diet drink. In a discussion question at the end of their reports, students are asked to consider whether their results are useful for predicting shelf lives of actual soft drinks. (For sure, the phosphate concentration and pH in a typical commercial drink are way below those in the experiment. But what about the effect of summer heat when drinks sit all day on a delivery truck?) Each reaction is initiated by adding solid aspartame (ca. 5 mg) to 10 mL of the desired buffer. After mixing, the solution is rapidly filtered into a tube and placed in the constant
H
H N
N
O
O
O OCH3
N O
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OCH3 H
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O
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Figure 1. Structures of (A) α-L-aspartyl-L-phenylalanine-1-methyl ester (aspartame) in the anion form and (B) 3-carboxymethyl-6-benzyl-2,5-dioxopiperazine (DKP).
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HO
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Scheme I. Proposed mechanism for conversion of aspartame to DKP.
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In the Laboratory Table 1. Kinetics Data from Two Student Groups kobs/min᎑1
Ea/(kJ mol᎑1)
Group
0.20 M Phosphate, pH 7.0, 40 ºC
0.20 M Phosphate, pH 7.0, 55 ºC
0.20 M Citrate, pH 7.0, 40 ºC
1
0.017 ± 0.003a
0.065 ± 0.010
0.0028 ± 0.0004
76 ± 13b
2
0.020 ± 0.003a
0.065 ± 0.007
0.0029 ± 0.0006
68 ± 11b
a
Uncertainties as 95% confidence limits.
b
Activation energies calculated with the phosphate data.
temperature bath. At suitable time intervals, 100 µL aliquots are withdrawn and transferred to chilled microfuge tubes containing 200 µL of phosphate buffer, 0.025 M pH 3.0. The combined effects of dilution, chilling, and decreases in both the buffer concentration and pH stop the reaction. As soon as possible, the diluted aliquots are separated on a 15-cm C-18 HPLC column using a mobile phase of 45兾55 methanol兾0.025 M phosphate, pH 3.0. First-order plots are obtained by graphing ln(At − A∞) versus time, where At and A∞ are the aspartame peak areas at times t and infinity, respectively. Hazards The methanol in the HPLC buffer may cause visual disturbances and possible blindness if inhaled or ingested. Skin contact can result in minor irritation or dryness (12). Results The first-order plots for the three reaction systems obtained by a typical student group are shown in Figure 2. Table 1 contains the rate constants (kobs) and activation energies (Ea) obtained by two student groups. The average activation energy of (71.9 ± 8.5) kJ兾mol compares favorably with the result of (58 ± 9) kJ兾mol obtained by Tsoubeli and Labuza
using 0.10 M phosphate, pH 7.0 (8). Bell and Wetzel obtained a rate constant of (0.00352 ± 0.00005) min᎑1 using 0.20 M phosphate, pH 7.0 at 25 ⬚C (11). Students use their observed activation energy and the rate constant at 40 ⬚C to calculate the value at 25 ⬚C. With the data in Table 1, the result is (0.00461 ± 0.00061) min᎑1, in reasonable agreement with Bell and Wetzel’s result. Conclusion The decomposition of aspartame is a popular experiment for physical and biophysical chemistry students, and it produces good results. With efficient usage of time, as specified in the Supplemental Material,W a single student group can complete the experiment in a five-hour laboratory period. Assuming that the course is operated on a rotation system with only one student group performing the experiment each week, the only major instrumentation requirement is a liquid chromatograph with a variable-wavelength UV detector. Students learn concepts of chemical kinetics and their relationship to the shelf-life of consumer products. W
Supplemental Material
Instructions for the students (including a prelab exercise) and notes for the instructor are available in this issue of JCE Online. Literature Cited
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Time / min Figure 2. Plots of ln(At − A∞) versus time at pH 7.0, where A is the HPLC peak area for aspartame: 0.20 M citrate at 40 ⬚C, diamond; 0.20 M phosphate at 40 ⬚C, square; 0.20 M phosphate at 55 ⬚C, circle.
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1. Delaney, Michael F.; Pasko, Kathleen M.; Mauro, David M.; Gsell, Diane S.; Korologos, Philip C.; Morawski, John; Krowlikowski, Linda J.; Warren, F. Vincent, Jr. J. Chem. Educ. 1985, 62, 618–620. 2. Conklin, Alfred R. J. Chem. Educ. 1987, 64, 1065–1066. 3. McDevitt, Valerie L.; Rodíguez, Alejandra; Williams, Kathryn R. J. Chem. Educ. 1998, 75, 625–629. 4. Bergen, H. Robert, III; Benson, Linda M.; Naylor, Stephen. J. Chem. Educ. 2000, 77, 1325–1326. 5. Lindeberg, Gunnar. J. Chem. Educ. 1987, 64, 1062–1064. 6. Stein, Paul J. J. Chem. Educ. 1997, 74, 1112–1113. 7. Stamp, Jeffrey Allen. Ph.D. Thesis, University of Minnesota, Minneapolis, MN, 1990. 8. Tsoubeli, Menexia N.; Labuza, Theodore P. J. Food Sci. 1991, 56, 1671–1675. 9. Prudel, M.; Davídková, E.; Davídek, J.; Kmínel, M. J. Food Sci. 1986, 51, 1393–1397. 10. Prankerd, R. J.; Stone, H. W.; Sloan, K. B.; Perrin, J. H. Int. J. Pharm. 1992, 88, 189–199. 11. Bell, Leonard N.; Wetzel, Clinton R. J. Agric. Food Chem. 1995, 43, 2608–2612. 12. EHS Services, Inc. Home Page. http://www.ehsservices.com (accessed Feb 2005).
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