In the Laboratory
Mercury-Free Analysis of Lead in Drinking Water by Anodic Stripping Square Wave Voltammetry
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Jeremy P. Wilburn, Kyle L. Brown, and David E. Cliffel* Department of Chemistry, Vanderbilt University, Nashville, TN 37235; *
[email protected] More and more smaller academic institutions are incorporating electrochemistry into their lab curriculum owing to the commercial availability of low-cost, PC-controlled electrochemical instrumentation. The development of experiments that use noble-metal electrodes (easily available if not supplied with the instrument package) and nontoxic materials could provide an added bonus for more institutions to add electrochemistry to their instrumental analysis and quantitative analysis courses. The analysis of drinking water for lead, which has well-known adverse health effects, provides an instructive example of the use of analytical chemistry to monitor a common hazard of everyday life. The technique of stripping voltammetry is well established, dating back to the works of Rogers (1) and Shain (2) in the 1950s. While the technique was, from its inception, utilized for the quantitative analysis of minute quantities of metals, its history has been largely tied to that of the hanging mercury drop electrode. The technique was featured in teaching curricula even prior to the availability of computercontrolled instrumentation (3–7), but fell into disfavor owing to the use of mercury electrodes. With the availability of low-cost microprocessor-controlled potentiometric workstations, stripping voltammetry is re-emerging as an option for inclusion in the teaching curriculum as well (8–11). Analysis of lead in tap water by anodic stripping voltammetry (ASV) has been established, but in its prior form was not desirable for undergraduate labs owing to the use of toxic materials. The traditional technique utilized an elemental mercury electrode that, when biased at the proper potential, forms an amalgam with the lead present in solution (12). Recently, Goebel et al. (10) reported a modified version of the technique, using a graphite electrode. While this is an improvement over the use of a mercury electrode, the technique still employs a toxic mercury salt to create an Hg–Pb amalgam in solution, which is then deposited on the surface of the electrode. Bonfil et al. (13) have described a mercuryfree analysis for lead in solution using a gold rotating-disc electrode, where an Au–Pb adlayer is deposited and analyzed via anodic stripping square wave voltammetry. While this technique absolved the need for hazardous materials, it requires a specialized electrode setup not common in teaching labs. Cypress Systems, manufacturers of electrochemical instrumentation, have published on their Web site a similar experiment (14) featuring a stationary glassy carbon electrode (GCE). However, owing to the relative ease with which the GCE can be damaged during polishing (which must be done between every analysis), we have found that it is not the most desirable electrode for incorporation into student lab experiments. Furthermore, the Cypress experiment focuses on a range of lead content 10 times higher than the Environmental Protection Agency’s (EPA) limit and 100 times that typically encountered in drinking water and also recommends that Hg salts be added to improve data quality. 312
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Bioanalytical Systems, however, in its in-house journal, Current Separations, has made the case for the use of gold electrodes for ASV of lead (15). In addition to using a more robust electrode surface, their method allows the analysis of lead levels more typically encountered in public drinking water. We have adopted their research methodology into a teaching experiment used in quantitative analysis labs for the past three years that allows student evaluation of lead levels in tap water. Theory Square wave voltammetry (SWV) is a widely used technique owing to its enhanced sensitivity, dynamic background suppression, and overall versatility (16). Likewise, ASV is recognized as being a highly sensitive form of quantitative analysis for dilute samples since the analyte is preconcentrated on the surface of the electrode. When using a gold electrode, the preconcentration results in the formation of a Au–Pb adlayer on the surface of the working electrode (WE), as shown in eq 1: + − Pb2 + 2e
Pb(Au)
E dep = − 0.6 V (vs. Ag/AgCl, 3 M KCl)
(1)
After 3 minutes of deposition, the preconcentration is sufficiently complete and two seconds of “quiet time” are enabled, allowing the deposited film to achieve a more uniform composition. The stripping mode consists of ramping the potential from the deposition potential (᎑0.6 V) to +0.1 V, thus causing the reverse electrochemical reaction to occur. The current resulting from the SWV stripping is proportional to the quantity of lead present in the solution. Experimental The electrochemical workstation (model 660A), the electrodes, and the polishing kit are all used as supplied by CH Instruments (Austin, TX). The technique employs a threeelectrode setup consisting of a 2-mm gold disc WE (model CHI101), platinum wire counter electrode (model CHI115), and Ag兾AgCl (3 M KCl) reference electrode (model CHI111). The electrodes are placed in an undivided 15-mL glass cell (a small glass weighing bottle), and held in place with the cell stand supplied with the instrument. Prior to each analysis, all electrode surfaces are cleaned or polished. The WE is polished by using slurries freshly prepared from 1-, 0.5-, and 0.05-µm alumina (from polishing kit model CHI120), which is placed on a micro-polishing cloth (a flat glass surface, such as a microscope slide, would work as well). The electrode is polished by moving across the top of the slurry in a figure-eight motion for approximately 3 minutes, and then thoroughly rinsed with deionized (DI) water to remove any adherent alumina. The counter electrode is flame
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In the Laboratory
polished by passing it through a methane Bunsen burner flame until all color subsides or thoroughly rinsed with DI water. The reference electrode is cleaned by rinsing the body and frit with copious quantities of DI water. A 500-ppb lead reference standard is prepared daily by serial dilution of a 50-ppm solution, which is made by the students according to their calculations. The students prepare this solution by dissolving the calculated quantity of PbCl2 in a 500-mL or 1-L volumetric flask. A dilution is then performed by transferring 1-mL (using a class A transfer pipet) of the 50-ppm solution to a 100-mL volumetric flask and diluting to the mark. The entire experiment is performed under quiescent conditions without the need for degassing. Prior to the deposition step a preconditioning mode (+0.8 V for 5 s) is utilized to electrochemically polish the WE surface (14–16). Ten mL of the sample to be analyzed (tap water, DI water, or an unknown solution previously prepared by the instructor) is then transferred (via a class A transfer pipet) to the glass cell containing the electrodes. A 100-µL spike of a 1 M KCl solution serves as the supporting electrolyte. Standard addition and calibration curves are performed by spiking the sample with successive 100-µL aliquots of the 500-ppb lead solution, each resulting in an increase of ∼5 ppb lead. We use an Eppendorf micropipetter to make the additions, but a microsyringe would work equally well. The specific experimental parameters are listed in Table 1.
lead (ppb, x axis). Calibration curves are used to show the linear response of the instrument in the concentration range expected for the analysis. For our system, the peak current is automatically measured by the supplied software. For systems lacking this ability, the method suggested by Goebel et al. (10) can be used instead. Determination of lead in tap water is performed via the standard addition method (Figure 2). The voltammogram of
Hazards Lead chloride (PbCl2) is listed as a high health hazard for chronic exposure via ingestion or inhalation. Severe overexposure can result in death. Additionally, fire or excessive heat can lead to decomposition into lead and hydrogen chloride fumes. Results A series of overlaid voltammograms, using the common electrochemical convention where negative values indicate oxidative currents, and the corresponding calibration curve for lead additions to DI water are presented in Figure 1. The calibration curve is obtained by plotting the absolute value of the individual peak height (current, y axis) versus the added
Figure 1. A series of voltammograms (top), and their corresponding calibration curve (bottom). The series was made by successive additions (each equivalent to ~5 ppb) of lead to DI water.
Table 1. Experimental Parameters Pre condit ioning M ode Pot e nt ial
+0.8 V
Time
5s
De pos it ion M ode Pot e nt ial
᎑0.6 V
Time
180 s
St ripping M ode Init ial pot e nt ial
᎑0.6 V
Final pot e nt ial
+0.1V
St e p incre me nt
0.004 V
St e p amplit ude
0.025 V
St e p f re que ncy
15 H z
Se ns it iv it y
5 x 10−6 A /V
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Figure 2. Standard addition plot for the analysis of lead in a sample of tap water (data from Table 2, group 5).
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In the Laboratory Table 2. Student Data for the Addition of Lead to Tap Water Current/(10᎑7 A) for Each Group
Pb Added/ ppb
1
2
3
4
5
6
7
00
3.04
2.52
4.17
6.18
5.00
3.57
4.58
05
3.47
3.50
6.33
7.99
6.48
4.44
5.60
10
4.93
4.86
7.62
9.85
7.53
5.62
7.16
15
---
5.68
9.24
---
8.88
6.50
8.41
20
---
---
---
---
---
7.14
---
Intercept
15
12
13
17
20
20
17
R2
0.911
0.992
0.989
0.999
0.996
0.991
0.995
Table 3. Student Data for the Addition of Lead to a 10-ppb Lead Solution Current/(10᎑6 A) for Each Group
Pb Added/ ppb
1
2
3
4
5
6
00
1.597
0.301
0.936
0.529
0.855
0.692
05
2.098
0.472
1.404
0.877
1.294
1.081
10
2.928
0.580
1.923
1.040
1.662
1.420
15
---
---
2.551
---
---
1.833
Intercept
12
11
8.4
9.2
11
9.2
R2
0.980
0.983
0.995
0.978
0.997
0.999
the unknown sample is initially taken after adding the supporting electrolyte and then after each of the three 100-µL spikes of the 500-ppb lead standard. The lead concentration is then determined from a plot of peak current versus concentration of added lead (as before). Once a best fit line is made, extrapolation is made backward to its x axis intercept (a negative value); taking the absolute value of that intercept will yield the lead concentration of the initial sample. The difference in current (one order of magnitude) between the calibration curve and the tap water sample can be related to interfering contributions in the sample matrix (17). This is an ideal opportunity to illustrate the deviations that can occur between signal兾concentration ratios observed in the matrix of a standard and in that of an actual sample and to reinforce the need for the method of standard addition in such cases. The data for the entire class can be seen in Table 2. The analysis results in an average lead content of 16.2 (±3.1) ppb. For the purposes of validating the results from the experiment, the students also perform a standard addition determination of an unknown solution prepared by the lab TA. Class data for such an experiment are presented in Table 3. The unknown solution was prepared at 10-ppb concentration and the average of the experimental determinations was 10.1 (±1.3) ppb. Conclusion This experiment combines many essential elements of experimental analytical chemistry; the ppm/ppb concept, standard preparation by serial dilution, instrument calibration, matrix effects, and unknown determination by standard addition. Finally, since the adverse effects of lead contamination to human health are well known, the experiment allows students to perform an experiment that can evaluate a risk factor in their own lives. 314
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Supplemental Material
Instructions for the students and notes for the instructor are available in this issue of JCE Online. Literature Cited 1. Lord, S. S., Jr.; O’Neill, R. C.; Rogers, L. B. J. Chem. Educ. 1952, 24, 209–213. 2. DeMars, R. D.; Shain, I. J. Chem. Educ. 1957, 29, 1825–1827. 3. Ellis, W. D. J. Chem. Educ. 1973, 50, A131, A35–A37, A40, A42–A47. 4. Deanhardt, M. L.; Dillard, J. W.; Hanck, K. W.; Switzer, W. L. J. Chem. Educ. 1977, 54, 55–57. 5. Stock, J. T. J. Chem. Educ. 1980, 57, A125–A26, A28, A30, A32–A34. 6. Wang, J. J. Chem. Educ. 1983, 60, 1074–1075. 7. Pomeroy, R. S.; Denton, M. B.; Armstrong, N. R. J. Chem. Educ. 1989, 66, 877–880. 8. Goscinska, T. J. Chem. Educ. 1998, 75, 1038–1039. 9. John, R.; Lord, D. J. Chem. Educ. 1999, 76, 1256–1258. 10. Goebel, A.; Vos, T.; Louwagie, A.; Lundbohm, L.; Brown, J. H. J. Chem. Educ. 2004, 81, 214–217. 11. Collado-Sanchez, C.; Hernandez-Brito, J. J.; Perez-Pena, J.; Torres-Padron, M. E.; Gelado-Caballero, M. D. J. Chem. Educ. 2005, 82, 271–273. 12. Deanhardt, M. L.; Dillard, J. W.; Hanck, K. W.; Switzer, W. L. J. Chem. Educ. 1977, 54, 55–57. 13. Bonfil, Y.; Brand, M.; Kirow-Eisner, E. Anal. Chem. Acta 2002, 464, 99–114. 14. Cypress Systems. http://www.cypresssystems.com/Experiments/ leadtap.html (accessed Oct 2006). 15. Jayaratna, H. G. Current Separations 1994, 12, 173–176. 16. Bard, A. J.; Faulkner, L. R. Electrochemical Methods; John Wiley & Sons: New York, 2000. 17. Harvey, D. Modern Analytical Chemistry; McGraw–Hill: Boston, 2000; p 110.
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