In the Laboratory
Troubleshooting 101: An Instrumental Analysis Experiment Joseph E. Vitt Department of Chemistry, University of South Dakota, Vermillion, SD 57069;
[email protected] Troubleshooting is an important skill that is difficult for students to learn in a traditional teaching lab, mainly as a result of time constraints. In an instrumental analysis course, there are simply too many instruments and techniques to learn during the course of a semester to allow much time for correcting errors. Thus, instructors diligently make certain that all instruments are in good working order and that the experiments will yield good results in the time allotted. But mistakes, of course, are often great learning opportunities (1). For one thing, mistakes are more memorable than easy successes, and critical thinking is not usually required if the first attempt is successful. Real learning occurs when our faulty understanding is confronted with evidence to the contrary, and we are forced to revise our understanding (2). Troubleshooting involves a systematic process of trial and error usually performed by making educated guesses about the source of the problem. Analytical chemists in industry are expected to be skilled at solving problems (3–5), and troubleshooting a faulty procedure, poor analytical result, or an instrument malfunction are important parts of that. Teachers of analytical chemistry have a long history of expressing dissatisfaction with the lack of problem-solving in the curriculum and proposing various solutions (3). Problem-based learning (PBL) has been emphasized in recent years to enable students to learn the process of analytical chemistry in addition to the traditional content (6–10). Many of the best published examples involve revising the entire curriculum around a central theme or analytical sample (11–17). Examples include semester-long cooperative-learning projects (11), courses designed around analysis of aquarium water (12), natural water systems (13–15), or urban air quality (16), and a course modeled on industrial laboratory projects (17). This experiment provides instructors a chance to implement PBL in one experiment to enhance a more traditional course or as part of a series of PBL labs. In this experiment students are asked to troubleshoot an analytical method for ethanol previously published in this Journal (18) that does not work well as written. It has been used in an instrumental analysis course for several years and takes one to two lab periods to complete. The instrumental analysis course includes third- and fourth-year chemistry majors and has a prerequisite of one semester of analytical chemistry. Ideally,
students will have experience with UV–vis spectroscopy prior to performing this experiment. The goal of this experiment is to help students develop their troubleshooting skills. Most instructors have probably observed students handing in lab reports with completely meaningless results due to a major error or misconception while performing a traditional lab. The major advantage of this experiment for my students is that it forces them to critically evaluate their experimental results during lab and to make decisions based on their understanding of their results—in other words, to “think on their feet”. To troubleshoot this experiment, students need to have a good understanding of the fundamentals of absorbance spectroscopy and be able to evaluate results and modify experimental conditions in a strategic way to determine the source of the problem. Because the stoichiometry of the reaction is fairly complicated, and involves a decrease in absorbance with addition of analyte, it is usually relatively difficult for students to determine the source of the error without significant thought and experimental work. If students do not evaluate and understand their results, they simply will not know when they have successfully completed the experiment. Procedure The experiment involves the analysis of ethanol by its reaction with dichromate, which is the chemical reaction used in the original Breathalyzer for the analytical determination of ethanol in breath: 3C2H5OH 2Cr2O72 16H
3CH3CO2H 4Cr3 11H2O
The subsequent decrease in concentration of dichromate is measured spectrophotometrically. The reaction conditions given in the original procedure (18) and an example of a successful student procedure are given in Table 1. Students are given the reference for the original procedure and instructed to determine whether the experiment works as written using the instrumentation available to them (HP diode array spectrophotometer or Spec 20). If the experiment does not work well,
Table 1. Solution Preparation and Expected Results Procedure
Solution Preparation
Expected Concentration
Dichromatea/mL
Ethanol
Water/mL
Dichromate/(10–4 M)
Blank
10
0
1
9.21
Original
10
1 mL of 5.31 x 10–4 M
0
8.86
0
5.96
Student aThe
1660
10
dichromate concentration is 1.013 x
1 mL of 5.31 x 10–3
10–3
M
M.
Journal of Chemical Education • Vol. 85 No. 12 December 2008 • www.JCE.DivCHED.org • © Division of Chemical Education
In the Laboratory 0.30
blank original student
0.25 0.20
Absorbance
they are instructed to devise a revised analytical procedure that could be used to determine the ethanol concentration in a sample and explain why the original procedure did not work as written. The experiment can be done in groups or individually, but it is a particularly good group experiment in that it requires students to discuss results and plan experiments based on those results.
0.15 0.10 0.05
Hazards
0.00
Potassium dichromate is a strong oxidant and carcinogen. Also, 50% sulfuric acid is used to dissolve the samples. Appropriate eye protection and protective gloves should be worn when preparing and handling solutions. Student Results Typical student results for the experiment are shown in Figure 1 and Table 2. The absorbance spectra obtained using the original procedure are indistinguishable for the solutions with (Figure 1, original) and without (Figure 1, blank) the addition of ethanol. The theoretical change in dichromate concentration is 4% for the original procedure, which might, in many cases, be easily detected spectrophotometrically. However, because of the relatively low molar absorptivity for dichromate (294 M–1 cm–1 based on student results), the theoretical change in absorbance is only 0.01. Given the limitations of the solution preparation, it is not unusual to get results that give a negative concentration of ethanol using the original procedure (as was the case with the results listed in Table 2), since the change in absorbance is smaller than the uncertainty of the experiment. Students often attempt to improve the procedure by focusing on the solution preparation without changing the concentrations, but eventually realize that the change in concentration of dichromate is too low given the sensitivity of the instrument. This provides a real lesson in the limitations of the instrument and absorbance measurements, making the earlier classroom discussion of figures of merit for instruments more meaningful. For example, the relatively poor sensitivity is due to the low molar absorptivity of dichromate. The spectrum obtained by students after revising the procedure to use a larger concentration of ethanol is shown in Figure 1. The decrease in absorbance under these conditions was easily measured, and the value calculated for the ethanol concentration from the spectrophotometric method was in good agreement with the value based on the volume of ethanol added (3.0% relative error). These values are listed for comparison in Table 2.
0.05 400
450
500
550
600
650
700
750
800
Wavelength / nm Figure 1. Absorbance spectra for dichromate: both the original procedure and the student procedure have ethanol added. See text and Table 1 for solution details.
Other approaches to troubleshooting the procedure are also possible. One solution my students have used is to add multiple aliquots of ethanol until the percent error is acceptable. Another solution might be to lower the concentration of dichromate in the procedure, thus getting a larger relative change in absorbance upon addition of ethanol. Although my students have not attempted this correction yet (it was suggested by a helpful reviewer of this manuscript), it could probably be elicited by requiring students to devise a procedure that would work for the ethanol concentration stated in the original procedure. This would have the benefit of helping students to think more deeply about the constraints of a real-world sample. Since the absolute decrease in absorbance would still be very small in this case, students might also be forced to consider other figures of merit. For example, they could try acquiring the signal for a longer time to improve the signal-to-noise and see whether that enables the spectrophotometer to measure the small change in absorbance. Conclusion Troubleshooting an analytical procedure is an important but often time consuming task. In the experiment presented here, students can successfully troubleshoot an experiment in one to two lab periods. Alternatively, faulty experimental procedures could be devised for students to modify, but using an experiment that has been published raises the interest level, making the exercise seem more “real world”, and teaches a valuable lesson about the scientific literature. Some instructors may wish to give students an unknown ethanol solution in addition
Table 2. Student Results before and after Troubleshooting Parameter
Original Procedure
Student Procedure
Absorbance for blank
0.271
0.271
Absorbance after addition of ethanol
0.273
0.179
Concentration of ethanol from solution preparation/M
5.31 x 10–4
5.31 x 10–3
Concentration of ethanol from spectrophotometry/M
–1.16 x 10–4
5.15 x 10–3
120
3.0
Relative error (%)
© Division of Chemical Education • www.JCE.DivCHED.org • Vol. 85 No. 12 December 2008 • Journal of Chemical Education
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In the Laboratory
to having them demonstrate a successful procedure for the analysis. Also, the experiment can be used to introduce students to method validation, which is critical for analytical methods in industrial labs. Students can then be required to write their reports in the form of a standard method for ethanol. Acknowledgments The author thanks the students who participated in this experiment, especially those whose data were used, including Shawn Bartel, Yali Yang, James Bower, and David Ring. Literature Cited Wenzel, T. J. Anal. Chem. 2002, 74, 439A–440A. Shiland, T. W. J. Chem. Educ. 1999, 76, 107–109. Woodget, B. W. Anal. Chem. 2003, 75, 307A–310A. DePalma, R. A.; Ullman, A. H. J. Chem. Educ. 1991, 68, 383–384. 5. Thorpe, T. M.; Ullman, A. H. Anal. Chem. 1996, 68, 477A– 480A. 6. Kuwana, T. Curricular Developments in the Analytical Sciences: A Report from NSF Workshops. National Science Foundation: Washington, DC, 1998. 7. Mabrouk, P. A. J. Chem. Educ. 1998, 75, 527–529.
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8. 9. 10. 11. 12. 13. 14. 15. 16. 17. 18.
Wenzel, T. J. Anal. Chem. 1999, 71, 693A–695A. Wenzel, T. J. Anal. Chem. 2001, 73, 501A–502A. Mabrouk, P. A. Anal. Chem. 2002, 74, 269A–274A. Wenzel, T. J. Anal. Chem. 1995, 67, 470A–475A. Hughes, K. D. Anal. Chem. 1993, 65, 883A–889A. Werner. T. C.; Tobiessen, P.; Lou, K. Anal. Chem. 2001, 73, 84A–87A. Arnold, R. J. J. Chem. Educ. 2003, 80, 58–60. Adami, G. J. Chem. Educ. 2006, 83, 253–256. Hope, W. W.; Johnson, L. P. Anal. Chem. 2000, 72, 460A– 467A. Wilson, G. S.; Anderson, M. R.; Lunte, C. E. Anal. Chem. 1999, 71, 677A–681A. Timmer, W. C. J. Chem. Educ. 1986, 63, 897–898.
Supporting JCE Online Material
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Journal of Chemical Education • Vol. 85 No. 12 December 2008 • www.JCE.DivCHED.org • © Division of Chemical Education
