In the Laboratory
Salicylate Detection by Complexation with Iron(III) and Optical Absorbance Spectroscopy An Undergraduate Quantitative Analysis Experiment Jeremy T. Mitchell-Koch* Department of Chemistry, Emporia State University, Emporia, KS 66801; *
[email protected] Kendra R. Reid and Mark E. Meyerhoff Department of Chemistry, University of Michigan, Ann Arbor, MI 48109
The analysis of common over-the-counter medications has become a popular approach in undergraduate analytical chemistry laboratories (1–4). The connection of the students’ course work to the “real” world is exciting and often leads to an increased interest in chemical analysis. This idea has been at the heart of the quantitative analysis laboratory courses at both Emporia State University and the University of Michigan, and our experiences led us to believe that our approach improves student learning and student enthusiasm for carrying out given laboratory experiments. Spectroscopy is a critical tool in the analysis of different molecules and is incorporated in the secondyear analytical chemistry laboratories. This experiment describes the quantitation of salicylate in liquid face wash using visible spectrophotometry. A similar colorimetric method published in this Journal described the detection of salicylate in urine for nursing students (5). Here, we expand the application of visible spectroscopy to determine salicylate in commercial face wash products and also to investigate the nature of the colored product formed during the analysis. There are several uses for salicylate and it is included in many everyday products. Salicylic acid is the major metabolite of aspirin and is commonly found in medications that treat acne, warts, and other similar ailments. Owing to the many medical applications of salicylic acid, a number of methods for its quantitation have been developed, including titration (6), gas–liquid chromatography (7), ultraviolet spectroscopy (8), and fluorescence spectroscopy (9, 10). The most widely used method in clinical laboratories, however, employs visible spectrophotometry where excess iron(III) is added in an acidic solution to form a highly colored species (11). A version of the visible spectrophotometry method is applied to quantitate salicylate in a commercial product and also in an instructor-prepared unknown solution. The Beer–Lambert law is followed over the concentration range of interest (10–100 mM), and a linear calibration curve is prepared. For the analysis of the commercial face wash, we collected and posted all trials for the entire class and asked students to conduct a t test comparing their data to that of the class. In addition, students compared their results with the salicylic acid value provided by the manufacturer. This requires students to practice unit conversions from amount concentration to mass concentration expressed as a percentage. Typical student results for this analysis compare favorably to the salicylic acid level that the manufacturer claims. A recent class at Emporia State University (8 students, 24 analyses) determined an average salicylic acid concentration of (1.01 ± 0.09)% compared to 1% listed by the manufacturer (only one significant figure provided). While these results generally confirm the accuracy of the analysis, synthetic unknown solutions 1658
are also used to further assess the students’ analytical skills. In these experiments all students generally determine the unknown concentration to within 5% of the expected value. These results indicate that students benefit from the opportunity the experiment provides to learn about a critical analytical tool. In the second part of the experiment, the method of continuous variation ( Job’s method) is used to investigate the stoichiometry of the iron–salicylate coordination complex formed. Varying volumes of 10 mM solutions of sodium salicylate and iron(III) nitrate are mixed and diluted to prepare eleven test solutions (see Table 1 in the online material). The total combined volume of the two solutions remains constant (1 mL), and thus only the mole ratio of iron to salicylate changes in these solutions. The test solution that exhibits the largest absorbance at λmax indicates the mole ratio of the predominant complex formed in the reaction. This maximum is most easily demonstrated in a plot of absorbance versus mole fraction of iron(III). Experimental Procedure Spectrophotometric Determination of Salicylate in Acne Medication A stock solution of approximately 100 mM salicylate is prepared from its dried sodium salt and is then appropriately diluted to yield five other standards for the calibration curve. To 100 μL of each standard, triplicate samples of face wash, and the synthetic unknown, 10.00 mL of acidic iron(III) nitrate solution is added. The absorbance of these solutions is then measured in plastic cuvettes. This step may be modified depending on available instrumentation. In one class, each pair of students had access to an Ocean Optics absorbance instrument that provided the entire spectrum for each sample. This allows for simple determination of λmax and the absorbance values at λmax. In another laboratory setting, one scanning spectrophotometer, and several single wavelength “spec 20” units were available. In this situation, each pair of students collected a spectrum of one standard to determine λmax. Then the spec 20 units were set to this wavelength and the absorbance of each calibrant and test solution was determined. If only single wavelength instruments are available, instructors may provide students with the λmax value to expedite the measurement process (~535 nm). From the absorbance data collected for the standards, a linear calibration curve is constructed (Figure 1), and the concentration of each unknown is determined using a least-squares fit line derived from the calibration data. The absorbance data are plotted against the undiluted standard concentrations to simplify quantitation of salicylate in each unknown. Since all standards and unknowns
Journal of Chemical Education • Vol. 85 No. 12 December 2008 • www.JCE.DivCHED.org • © Division of Chemical Education
In the Laboratory
are diluted in the same manner, the salicylate concentrations in the unknown solutions can be extracted directly from the calibration curve without any dilution calculations. This method is valid for concentration determination, but the plot will underestimate the extinction coefficient for the iron–salicylate complex by approximately two orders of magnitude. The experiment may be modified to require a plot of absorbance versus actual salicylate concentration (after dilution in the reaction mixture) to properly determine the extinction coefficient. Determination of Reaction Stoichiometry: An Application of the Method of Continuous Variation Various volumes of 10 mM iron(III) nitrate and 10 mM sodium salicylate are mixed in separate test tubes in ratios shown in Table 1 (found in the online material). The total volume of these two solutions is a constant 1.00 mL, keeping the total amount of reagents the same. Dilute nitric acid (4.00 mL, ~60 mM) is added to each mixture, and the absorbance at λmax is recorded for each solution. A typical plot of absorbance versus the mole fraction of iron(III) is shown in Figure 2. An apparent local maximum around 0.250 mole fraction (a 1:3 iron:salicylate complex) is a typical result, which may represent a minor complex, but the overall maximum in this plot at 0.500 mole fraction (a 1:1 iron:salicylate complex) indicates the predominant stoichiometry of the reaction.
Absorbance at 535 nm
1.2
Nitric acid is corrosive, and iron(III) nitrate is an oxidizer: appropriate safety precautions should be taken when handling these chemicals. Summary The experiment provides students practical experience with Beer–Lambert law for the spectroscopic quantitation of molecules in aqueous solution. In addition, spectroscopy is used to investigate the nature of a chemical reaction and determine the stoichiometry of the absorbing species. The experiment can be completed in a single three-hour laboratory session but may be divided to provide students ample time to analyze their data as it is collected. The reagents are common, inexpensive, and require no special safety precautions. The experiment can be easily adapted to a wide variety of spectrophotometers and therefore is simple to adopt into any second-year quantitative analysis course. This experiment is relatively simple to perform and helps students gain an understanding of optical spectrophotometry and equilibrium. If appropriate analytical techniques are used, students encounter little difficulty in obtaining a linear calibration curve and determining the predominant complex stoichiometry. Student results demonstrate excellent accuracy (vide supra), and this enhances the enjoyment of the activity for both students and instructors. Acknowledgments
1.0
The authors acknowledge the University of Michigan for funding and all the students at U of M and ESU who tested this experiment.
0.8 0.6
Literature Cited
0.4 0.2 0.0
0
20
40
60
80
Salicylate Concentration / (mmol/L) Figure 1. Typical calibration curve for salicylate.
1. 2. 3. 4. 5.
6. 7. 8. 9. 10. 11.
0.6
Absorbance at 535 nm
Hazards
0.5
0.4
Ferguson, G. K. J. Chem. Educ. 1998, 75, 467–469. Hein, J.; Jeannot, M. J. Chem. Educ. 2001, 78, 224–225. Simonson, L. A. J. Chem. Educ. 2001, 78, 1387. Yang, S.-P.; Tsai, R.-Y. J. Chem. Educ. 2006, 83, 906–909. Cavanaugh, M. A.; Bambenek, M. A. J. Chem. Educ. 1978, 55, 464. Lane, S. R.; Stewart, J. T. J. Chem. Educ. 1974, 51, 588–589. Battezzati, A.; Fiorillo, G.; Spadafranca, A.; Bertoli, S.; Testolin, G. Anal. Biochem. 2006, 354, 274–278. Rogic, D. J. Mol. Struc. 1993, 294, 255–258. Lange, W. E.; Bell, S. A. J. Pharm. Sci. 1966, 55, 386–389. Saltzman, A. J. Biol. Chem. 1948, 174, 399–404. Annino, J. S.; Giese, R. W. Clinical Chemistry: Principles and Procedures, 4th ed.; Little, Brown and Co.: Boston, 1976; pp 355–357.
Supporting JCE Online Material
0.3
http://www.jce.divched.org/Journal/Issues/2008/Dec/1658.html Abstract and keywords
0.2
0.1 0.0
Full text (PDF) with links to cited JCE articles 0.2
0.4
0.6
0.8
Mole Fraction Fe(III) Figure 2. Typical Job Plot for the iron-salicylate complex.
1.0
Supplement
Student handouts
Instructor notes
© Division of Chemical Education • www.JCE.DivCHED.org • Vol. 85 No. 12 December 2008 • Journal of Chemical Education
1659
