In the Laboratory
Waste Treatment in the Undergraduate Laboratory: Let the Students Do It! John J. Nash,* Jeanne A. R. Meyer, and Susan C. Nurrenbern Department of Chemistry, Purdue University, West Lafayette, IN 47907 There are numerous reports in this Journal that describe procedures for treating various types of wastes generated in research labs as well as in undergraduate teaching labs (1–11). However, relatively few reports describe direct student involvement in the waste treatment process (12–18). Including well-designed waste treatment “experiments” as natural addenda to laboratory experiments that generate hazardous waste could provide students additional opportunities to apply chemistry to “real” problems. Moreover, such waste treatmenttype experiments might be used to stimulate discussions regarding the hazards and environmental issues associated with the disposal of chemical wastes at both the academic and industrial levels. We were prompted by the paper of Schneider and Wiskamp (18) to present the details of a waste treatment experiment that we have recently incorporated into our large (ca. 2400 students per year), second-semester general chemistry course. We are using a modified version of the copper penny experiment reported by Vanselow and Forrester (19). Students dissolve a copper-clad penny1 in nitric acid, complex the copper with ammonia, and determine the copper content in the penny spectrophotometrically using a separately measured calibration curve. During the course of this experiment, students save all chemical waste in a covered beaker. At the end of the experiment, each group of four students has collected about 300 mL of waste (a mixture of aqueous and solid wastes) that we estimate2 contains about 3800 ppm of copper and 9300 ppm of zinc. Because ammonia is used in the experiment, the pH of the waste is approximately 9. This produces a waste mixture containing a large amount of white, solid zinc hydroxide and a dark blue solution that presumably contains various copper and zinc complex ion species. The concentrations of zinc and copper in the waste mixture are much too high to be disposed of in the sanitary sewerage system. In fact, at Purdue University we are not allowed to discard aqueous wastes that contain more than 2.5 ppm of copper or more than 4.2 ppm zinc. Because of these restrictions, we originally collected the waste and disposed of it according to the guidelines set by the Radiological and Environmental Management Department at Purdue University. However, the students now chemically treat the waste they generate to remove the copper and zinc as a separate “experiment”. This approach has several advantages. For example, the students employ much of the chemistry (e.g., acid–base equilibria, solubility equilibria, and complex ion equilibria) that they have learned in class to solve a “real” chemical problem. Moreover, the volume of hazardous waste that we must dispose of is reduced by several orders of magnitude. Finally, the experiment heightens the student’s awareness of the potential environmental im*Corresponding author.
pact associated with the waste they have generated (see below). Chemistry Involved Because ammonia is added to most of the solutions prepared in the copper-penny experiment, the waste is slightly basic (pH about 9) and the copper is present primarily as the royal-blue tetraammine complex ion, Cu(NH3)42+. However, the zinc is probably present not only as the colorless tetraammine complex ion, Zn(NH3) 42+, but also as the sparingly soluble hydroxide salt, Zn(OH)2 (K sp = 4.5␣ ×␣ 10{17 ).3 All of these species are relatively stable. The first step involves neutralizing the ammonia in the mixture with hydrochloric acid to free the Cu2+ and Zn2+ ions: HCl(aq) + NH3(aq) → NH4+(aq) + Cl{(aq)
(1)
The addition of a sufficient amount of acid destroys the complex ions of both copper and zinc by shifting the following complex ion equilibria completely to the left: Cu2+(aq) + 4 NH3(aq) 2+
Zn (aq) + 4 NH3(aq)
Cu(NH3)42+
(2)
2+ 3 4
(3)
Zn(NH )
Moreover, the addition of acid also destroys any zinc hydroxide (or copper hydroxide) in the waste mixture by shifting the following equilibria to the right: Zn(OH)2(s) Cu(OH)2(s)
Zn2+(aq) + 2 OH{(aq) 2+
{
Cu (aq) + 2 OH (aq)
(4) (5)
At this point, the waste solution contains the aqueous zinc and copper ions. The zinc and copper are now precipitated as the sulfides by adding a solution of sodium sulfide: Cu2+(aq) + S2{(aq) → CuS (s)
(6)
Zn2+(aq) + S2{(aq) → ZnS (s)
(7)
Virtually all of the zinc and copper is precipitated from the solution as the sulfides because of the extremely low solubilities of these salts: Ksp (ZnS)␣ =␣ 1␣ ×␣ 10 {27 and K sp(CuS)␣ =␣ 6.7␣ ×␣ 10{42 . The sulfide precipitates appear to be colloidal and are difficult to remove by filtration. Therefore, we use a “cationic” (i.e., high positive charge density) polyelectrolyte to effect coagulation of the colloidal particles. We believe that the colloidal particles of CuS and ZnS are probably surrounded by an “electrical double layer” that prevents them from coagulating. This phenomenon is similar to that sometimes observed in gravimetric analyses (20). Moreover, it is known (20) that coagulation can sometimes be induced via the addition of electrolytes (or polyelectrolytes, in this case). The action of the polymer is probably related to its ability to reduce the electrical
Vol. 73 No. 12 December 1996 • Journal of Chemical Education
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In the Laboratory
double layer surrounding the colloidal CuS and ZnS particles, thereby allowing the particles to coagulate. Experimental Procedure CAUTION: Hydrogen sulfide gas is generated in this experiment; therefore, all operations should be carried out in a well-ventilated hood! Concentrated HCl and 6 M NaOH are corrosive. Gloves and eye protection should be worn at all times! The 300 mL of waste is divided into roughly equal portions and each pair of students performs the following procedure. Deionized water (100 mL) is added to 150 mL of waste mixture and concentrated HCl is added dropwise with stirring until any precipitate just dissolves and the solution is clear and light blue.4 Twenty-five milliliters of a 1.0 M Na2S solution is then added to precipitate the copper and zinc sulfides and the solution is stirred for an additional 2–3 min. CAUTION: Significant hydrogen sulfide gas will be generated! The pH of the mixture is then checked using “pHydrion A” pH paper (Micro Essential Laboratory Inc., Brooklyn, NY 11210).5 If necessary, 1.0␣ M Na2S is added dropwise to the mixture until the pH is ca. 6. The pH of the mixture is then adjusted to ca. 8 by dropwise addition of 6 M NaOH. The finely divided mixture of CuS and ZnS is then flocculated by adding 3 mL of a watersoluble organic polymer (Calgon Co., Pittsburgh, PA 15230, #WT-2640)6 and stirring the mixture for about 5 min. The mixture of solids is then removed by suction filtration7 and deposited in an appropriate waste container.8 Results If the students are careful during the filtration, the filtrate they obtain will be clear and colorless.9 The students then verify that the zinc has been removed from the filtrate, using zinc test strips (E. M. Science, Gibbstown, NJ 08027, #10038).10 To verify the absence of copper in the filtrate, the students add a small amount of concentrated ammonia to a portion of their filtrate. The absence of the dark blue copper–ammonia complex indicates the absence of copper. It should be noted that copper test strips (E. M. Science, #10003) are also available, but they are more expensive than the zinc test strips. In our case, one copper test strip is used to test one randomly chosen filtrate in each lab section.11 Students generally do not detect any zinc or copper in their filtrates. We independently analyzed several samples of filtrates obtained by our students by the inductively coupled plasma (ICP) (21) method12 and found that copper and zinc can, in fact, be reduced to about 0.2 and 1.2 ppm, respectively, using this procedure. After the students have completed their tests for zinc and copper, the filtrate is poured down the drain.13 Students have responded to this experiment surprisingly well, and many choose it as their favorite laboratory experiment of the semester! Their comments suggest that this experiment not only helps them to better understand the various aqueous equilibria involved in complex systems, but also provides them with a realworld example of applied chemistry. Moreover, students are particularly fascinated by the organic polymer used to flocculate the finely divided solids. Organic flocculents are used extensively in industrial waste treatment pro-
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cesses, and we typically devote some class time to a discussion of these applications. We believe that the procedure outlined here may have broad utility. For example, we are testing this methodology on aqueous wastes generated from our Qualitative Analysis and Electrochemical Cells experiments. Conclusion Our efforts to reduce the large volumes of hazardous chemical wastes generated in our undergraduate laboratories have resulted in a unique, pedagogically useful experience for our students. Clearly, the students understand and willingly accept responsibility for reducing the environmental impact associated with the chemical waste generated in their laboratory experiment and seem genuinely satisfied upon accomplishing this task. The following student comment reflects the ideas expressed in many of the laboratory reports for this experiment: “In other experiments, we simply disposed of our wastes in the waste jars provided with no thought of their fate...We learned that the principles of chemistry, when soundly practiced, can be of tremendous value in solving real-world problems.” Acknowledgments We are grateful to T. E. Van Dusen for providing the ICP measurements. We also thank the teaching assistants for implementing the procedure in the laboratories. Notes 1. Pennies minted in 1982 and thereafter are composed of a copper coating on a primarily zinc core. 2. We estimated these quantities based on the masses and volumes of the various reagents used in the experiment. 3. We recognize that the waste solution probably contains a variety of equilibria involving numerous copper and zinc species. We intentionally simplified the presentation of the chemistry to avoid these complexities. 4. We have observed that colloidal suspensions of elemental sulfur are produced if too much concentrated HCl is added to the solution. 5. This pH paper is used throughout the experimental procedure to monitor pH. 6. The polymer is a copolymer of dimethyldiallyl ammonium chloride and acrylamide. 7. It is possible that a filter aid (such as Celite) might be just as effective as the polymer for removing the precipitates. We have not, however, explored this possibility. 8. The disposal of the solid waste is handled by the Radiological and Environmental Management Department at Purdue University. 9. If the filtration is not done carefully, contamination of the filtrate with the dark brown solid will result. In this case, the filtration must be repeated. 10. The zinc test strips give the concentration of zinc in parts per million (ppm). 11. We generally find that the removal of copper from the waste mixture is more complete than the removal of zinc. 12. “ICP” is a type of emission spectroscopy that can be used to measure concentrations in the parts per million (ppm) to parts per billion (ppb) range of transition metal ions in aqueous solution. 13. All discharge should be done within the guidelines of any local and state discharge limitations.
Literature Cited 1. Bonner, W. D.; Masaki, K. J. Chem. Educ. 1930, 7, 616–617. 2. DeWitt, C. C. J. Chem. Educ. 1937, 14, 215–217.
Journal of Chemical Education • Vol. 73 No. 12 December 1996
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3. 4. 5. 6. 7. 8. 9. 10. 11. 12. 13. 14. 15. 16. 17. 18. 19. 20. 21.
Shirai, M.; Matsumoto, Y. J. Chem. Educ. 1977, 54, 609. Chamot, E. J. Chem. Educ. 1977, 54, 665. Bush, K. J.; Diehl, H. J. Chem. Educ. 1979, 56, 54–55. Armour, M. A.; Browne, L. M.; Weir, G. L. J. Chem. Educ. 1985, 62, A93–A95. Irving, N. M. J. Chem. Educ. 1986, 63, 1016. Rawat, J. P.; Kamoonpuri, S. I. M. J. Chem. Educ. 1986, 63, 537–538. Kauffman, G. B. J. Chem. Educ. 1988, 65, 375. Armour, M.-A. J. Chem. Educ. 1988, 65, A64–A68. Hubler-Blank, B.; Witt, M.; Roesky, H. W. J. Chem. Educ. 1993, 70, 408–409. Johnson, L. D. J. Chem. Educ. 1939, 16, 495–496. Pitt, M. J. J. Chem. Educ. 1980, 57, A261–A264. Chang, J. C.; Levine, S. P.; Simmons, M. S. J. Chem. Educ. 1986, 63, 640–643. Alvaro, M.; Espla, M.; Llinares, J.; Martinez-Manez, R.; Soto, J. J. Chem. Educ. 1993, 70, A129. Dhawale, S. W. J. Chem. Educ. 1993, 70, 395–397. McSwiney, H. D. J. Chem. Educ. 1994, 71, 329. Schneider, J.; Wiskamp, V. J. Chem. Educ. 1994, 71, 587–589, and references therein. Vanselow, C. H.; Forrester, S. R. J. Chem. Educ. 1993, 70, 1023–1024. Skoog, D. A.; West, D. M. Fundamentals of Analytical Chemistry, 4th ed.; Saunders: New York, 1982. Skoog, D. A. Principles of Instrumental Analysis, 3rd ed.; Saunders: New York, 1985.
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