In the Laboratory edited by
Cost-Effective Teacher
Harold H. Harris University of Missouri—St. Louis St. Louis, MO 63121
Small-Scale and Low-Cost Galvanic Cells
Per-Odd Eggen School Laboratory for Science and Technology, Norwegian University of Science and Technology, Trondheim, Norway Truls Grønneberg Department of Chemistry, University of Oslo, 0315 Oslo, Norway Lise Kvittingen* Department of Chemistry, Norwegian University of Science and Technology, 7491 Trondheim, Norway; *
[email protected] This article describes the construction of two small, simple, and illustrative galvanic cells that can be made by most students in any laboratory or most classrooms. These microscale constructions were developed during the preparation of courses in oxidation–reduction chemistry for secondary school teachers. Galvanic cells are traditionally presented in the form of the Daniell cell, both when used qualitatively and quantitatively. Cheaper and simpler varieties have been reported as well as adaptations to improve both comprehension and ease of demonstration (1–4). The two apparatuses we describe are shown in Figures 1 and 2. In both experiments students construct their own from inexpensive materials. We strongly believe that students should make their own devices when possible; however, materials should then be inexpensive and constructions quickly completed. Electrochemistry is, in general, considered to be difficult, and misconceptions are abundant and documented (5– 10). In our experience students have problems understanding voltage and current as well as the relation between these two concepts. This also applies when they study a Daniell cell. Digital and analog readings of voltage and current are not intuitively helpful either. We have therefore constructed two apparatuses in which these problems are avoided and instead lead to the simple (and understandable) conclusion, namely; that electric energy is generated (the glowing diode) from chemical reactions. The semiconducting property of the light diode can, of course, be exploited to identify at which electrode electrons are produced, thus where oxidation takes place, in these apparatuses (as elsewhere in small-scale electrochemistry). For students at higher levels voltmeters can be connected. As a first introduction to galvanic cells the classic Daniell cell, the Zn|Zn2+||Cu2+|Cu cell, is often presented. It contains “redundant” parts, namely the Cu metal and the Zn2+ solution, that often confuse students. Adding the “inverted horseshoe”, namely the salt bridge, incomprehension is often complete. We have therefore (in apparatuses 1 and 2) avoided the Zn2+ solution and the mysterious salt bridge and instead combined them in an inert salt solution physically held in place by floral foam. The Cu metal is replaced by an electrode of an “inert” material, which copper(II) ions cannot oxidize, for example, a silver wire, a graphite rod, or a pencil lead. On both these electrodes a copper deposit will easily be spotted. Each material has convenient properties; a silver wire is flexible and thus easy to handle, but a graphite rod is cheap www.JCE.DivCHED.org
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and available and also resembles the positive electrode of a common battery. Of course, if a silver wire or graphite rod is not available, a copper wire can be used. Which electrode to choose is therefore a consideration for the individual teacher. It is also worthwhile noting that the Daniell cell was neither one of the first voltaic cells made (11) nor does it resemble any battery or cell that the students are likely to meet outside the school laboratory, whereas our cells do. The chemistry involved in these experiments is not discussed in detail as this can be found in any general chemistry book (e.g., ref 12 ), as well as in the references cited above. Materials used to construct these two apparatuses are as follows: a culture plate (12 or 24 wells or small vials), disposable polyethene pipets (Beral pipets), small crocodile clips, graphite rod or pencil leads, galvanized nails or zinc strips, magnesium ribbon, silver wire of about 0.3-mm diameter. Smithers–Oasis Ideal floral foam (an open celled phenolic foam intended for live flowers) (13), light diodes (RS-components AS: RS Stock no 826-442), and saturated solutions of copper(II) sulfate and sodium sulfate. Safety goggles should be used during the experiment even though the risk of injury is very low. Visual Effect of One Cell: Energy from Copper(II) and Magnesium The first apparatus is shown in Figure 1. It produces a voltage and current sufficient to make the light diode glow. It is made as follows: 1. Shape, with (plastic) knife, a piece of the floral foam so that it fills at least half of a well (vial) and place it in the well. The piece of foam should be slightly taller than the depth of the well. 2. Push a magnesium ribbon (3–4 cm) into the floral foam as shown in Figure 1. The magnesium metal should not extend into the part of the floral foam that will be in contact with the copper salt solution, which will be added later; otherwise the magnesium will react directly with the copper salt solution. 3. Put a graphite rod (or silver wire) in the empty space of the vial. 4. With a pipet add inert salt solution (e.g., saturated sodium sulfate) corresponding to the salt bridge onto the floral foam until it cannot soak up any more.
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In the Laboratory 5. Connect a light diode to the magnesium ribbon and the graphite rod (or silver wire) with crocodile clips. Remember to put the longer “leg” of the diode onto the positive pole, that is, the graphite rod (or silver wire), otherwise no light will shine. (Alternatively hold the diode with your fingers onto the two electrodes after all chemicals have been added.) 6. Use a pipet to add copper(II) solution (e.g., saturated copper sulfate) to the vial as shown in Figure 1. Avoid spilling copper solution on top of the floral foam, otherwise it will react directly with the magnesium ribbon.
The diode will shine immediately but fade after a while (1– 24 hours), often because of magnesium depletion. The light diode must be of a type that requires low voltage and current and gives strong light (see above). The advantage of this setup is its simplicity: a metal and a metal solution apparently giving electric energy through chemical reactions.
Clearly the metal is oxidized, as corrosion of the magnesium ribbon can be verified by inspection, however, this is most likely from the direct reaction between magnesium and the aqueous solution held in place by the floral foam, and is not a result of the electric current passing through the diode. The detail of explanation is again a consideration for the individual teacher. It is also possible to compare this cell with a used flashlight battery (of the carbon zinc type normally used for heavy duty, e.g., an AAA Panasonic, heavy-duty battery) partly stripped of the case, if the teacher can provide one. The stripping of the case should however be left to the teacher and should only be conducted if he or she feels comfortable with this. In our cell (Figure 1) there are many possible reactions taking place especially with respect to what is being reduced; H2O to H2, O2 to OH−, or Cu2+ to Cu? To avoid these problematic side reactions with the reactive magnesium metal, we recommend constructing cells combining Zn strips (galvanized nails) and copper(II) solution, which are the standard combination in most chemistry books. In order to obtain a sufficient voltage for a light diode to glow a series of two or three cells is necessary. Visual Effect of a Small Battery: Energy from a Copper(II) and Zinc A complete battery is made of two to four cells. A twocell battery is shown in Figure 2. It is made as follows: 1. Shape two pieces of floral foam, so that each fills at least half (or most) of the area of the well and at the same time is slightly taller than the depth of a well.
Figure 1. A galvanic cell comprising a Mg ribbon (negative electrode) in a neutral salt solution (sodium sulfate) held in place by floral foam and an inert electrode (graphite rod, positive electrode) in copper sulfate solution. The electrodes are connected to a light diode.
2. Put one piece of foam in each of two neighboring wells as in Figure 2. (Some light diodes will require three cells in series in order to shine.) 3. Push a galvanized nail (or zinc strip) into the floral foam in the two wells. The nail should not extend into the part of the floral foam that will be in contact with the copper salt solution, which will be added later. 4. Twist one end of a silver wire around (or attach a graphite rod to) the nail in one well (right-hand well in Figure 2) and put the other end into the empty space in the adjacent well (left-hand well in Figure 2) where copper(II) salt solution will be added. 5. Put a silver wire (or graphite rod as shown in Figure 2) in the empty space in the first well (right-hand well in Figure 2). If a silver wire is not available, do not use a copper wire here, but a graphite rod, otherwise the copper layer deposited during the reaction cannot be spotted. 6. Connect one diode leg to the nail, which has no silver wire around (left-hand well), and the other leg to the graphite rod (or silver wire) in the other well (right-hand in Figure 2) for example using crocodile clips. As earlier mentioned, electrons can only flow in one direction in diodes, so connect the shorter “leg” to the nail.
Figure 2. Two galvanic cells in series. The negative electrode in each cell is a galvanized nail in neutral salt solution (sodium sulfate) held in place by floral foam. The positive electrode is a silver wire (left-hand well) and a graphite rod (right-hand well). Both positive electrodes are immersed in copper sulfate solution. A diode completes the circuit.
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7. Add inert salt solution, (e.g., saturated sodium sulfate solution) corresponding to the salt bridge, onto the floral foams until they cannot soak up any more. 8. Add copper(II) solution (e.g., saturated copper sulfate solution) into the empty spaces of the wells. Avoid spilling onto the top of the foam.
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In the Laboratory
The light diode will immediately start to shine, and from the semiconducting properties the students can infer that the electrons flow from the nail (negative pole) to the silver (graphite) electrode (positive pole), thus easily recognizing what is being reduced and what is being oxidized. Visual evidence of this is also easily obtained. The nails (or Zn strips) can be pulled out of the foam and inspected, revealing that the zinc has corroded, although this requires some time. (Short-cutting the circuit will, however, speed up this reaction.) At the silver wire (or graphite rod) a shiny copper layer will be deposited within 10 to 30 minutes. With this battery the diode will most likely glow until the next day. If not, it is normally only necessary to add a little copper(II) solution. More cells can, of course, be assembled in order to obtain a more powerful battery and voltmeters can be connected. We suggest the students study and compare drawings of classical dry cells, for example, Leclanché cell, found in most common general chemistry books including ref 12, with the ones the students have made. In common flashlight batteries it is not evident what is being reduced, in ours, however, the copper deposited onto the silver wire (or graphite rod) hopefully facilitates this understanding. Hazards Copper salts (chlorides and sulfates) may cause skin irritation as well as irritation of mucous membranes. Acute poisoning may occur, although such quantities normally lead to vomiting. Direct eye contact causes strong irritation. Magnesium may react with water to produce hydrogen, a flammable and explosive gas. The rest of the equipment (floral foam, graphite rod, silver wire, pipet, sodium sulfate, light diode, crocodile clips, wire and Zn-strips) are not hazardous.
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Acknowledgments We are grateful to Florin Banica, Ana Iona, Øyvind Mikkelsen, Knut Schrøder, and Brit Skaugrud for scientific advice. Literature Cited 1. Weinbert, N. L. J. Chem. Educ. 1972, 49, 120. 2. Thomson, S. Chemtrek, Small-Scale Experiments for General Chemistry; Prentice-Hall Inc.: New York, 1989; pp 300–314. 3. Slater, A.; Rayner-Canham, G. Microscale Chemistry Laboratory Manual; Addison-Wesley Pub. Lim.: Reading, MA, 1994; pp 151–176. 4. Ciardullo, C. V. Micro Action Chemistry; Flinn Scientific: Batavia, IL, 1992; Vol. 1, pp 75–77. 5. Özkaya, A. R. J. Chem. Educ. 2002, 79, 735–738. 6. Sanger, M. J.; Greenbowe, T. J. J. Res. Sci. Teach. 1997, 34, 377–398. 7. Birss, V. I.; Truax, D. R. J. Chem. Educ. 1990, 67, 403–408. 8. Garnett, P. J.; Treagust, D. F. J. Res. Sci. Teach. 1992, 29, 121– 142. 9. Sanger, M. J.; Greenbowe, T. J. J. Chem. Educ. 1997, 74, 819– 823. 10. Ogude, A. N.; Bradley, J. D. J. Chem. Educ. 1994, 71, 29–34. 11. Figuier, L. Les Merveilles de l’Electricité; Textes choisis (reprint of selected texts by Cardot, F. (1985), ISBN 2905821019, pp 480–500, Association pour l’histoire de l’électricité en France 1985, 47, rue de Monceau, 75008 Paris, originally from Les Merveilles de la Science 1867–1869. 12. Zumdahl, S. S.; Zumdahl, S. A. Chemistry, 5th ed.; Houghton Mifflin: Boston, 2000; pp 862–866. 13. Smithers-Oasis Worldwide Resource Center. http:// www.smithersoasis.com/res_techfacts.php (accessed May 2006).
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