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 Electrodes for “Standard” Reduction Potential Measurements Per-Odd Eggen and Lise Kvittingen* Department of Chemistry, Norwegian University of Science and Technology, 7491 Trondheim, Norway; *
[email protected] Truls Grønneberg Department of Chemistry, University of Oslo, 0315 Oslo, Norway
This article describes how to construct three simple and inexpensive electrodes: a hydrogen, a chlorine, and a copper electrode. These electrodes can be made by most students in any laboratory or most classrooms. We strongly believe in students making their own devices when possible; however, materials should then be inexpensive and constructions quickly completed. Central to most chemistry curricula is the measurement of reduction potentials. This is traditionally done using a reference scale and reference electrodes. Conventionally, the reference scale is the standard hydrogen scale; however, the corresponding reference electrode is not easily obtained. Actually a rigorous standard hydrogen electrode has even been termed a misrepresented concept (1). Hydrogen reference electrodes that can be made are considered inconvenient; therefore calomel or silver兾silver chloride electrodes are used instead. In a school setting it is crucial to use simple, inexpensive, and, if possible, home-made devices, which the above mentioned electrodes are not. We have therefore constructed a robust reference electrode, namely, a copper electrode. Simple copper reference electrodes have been reported earlier (2–4). The advantage of ours is that it is made from plastic and has no open ends (Figure 1). Thus leakage of solution is completely avoided, even when the electrode is unintentionally left horizontal on the bench, a situation quite likely when many students are at work. To our knowledge, no “home-made” hydrogen electrode as simple as ours has been reported before, although varieties exist (5–7). Ours is inexpensive, easily constructed, and gives measurements satisfactory for the intended purpose. Our hydrogen electrode incorporates platinum wires in disposable plastic pipets and as a bonus a chlorine electrode is obtained. This hydrogen electrode can be used to determine the “standard” potentials of various other electrodes. The hydrogen electrode is less stable than the copper electrode we described above. Therefore the preferred use of the hydrogen electrode is for calibrating results obtained from the copper reference electrode. This is done by measuring the potential between the copper reference electrode and the hydrogen electrode, assigning the value 0.0 V to the half cell potential of the latter. Of course the electrodes we describe here and the measuring of “standard” reduction potentials by the use of an inexpensive multimeter do not comply with rigorous standards for such measurements—especially the use of a multimeter and not a potentiometer—and may allow www.JCE.DivCHED.org
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small currents to flow, which then could polarize the electrodes. However the intended use of these electrodes is in secondary school and in first-year course at college or university where such rigor is not required. In electrochemistry misconceptions flourish (8–11). We believe that measuring cell potentials relative to a copper electrode instead of a hydrogen electrode is one way of approaching these. For example students will better understand the relative nature of cell potentials and they will also be less prone to believe that cell potentials can be measured directly. The chemistry involved in the experiments that can be done with these electrodes is not discussed here as this can be found in any general chemistry book. Materials used to construct these three electrodes are disposable, small polyethene pipets (Elkay Eireann, mini, stem-size 0.5 mL), Pt wires1 for the hydrogen and chlorine electrodes (e.g., 0.3 mm diameter × 4 cm, ∼$5 each), and copper wire for the copper reference electrode, as well as vials for the electrolytic cells. Various metals and their salt solutions are needed to make half-cells for reduction potential measurements with the copper reference
Figure 1. A copper reference electrode made from a plastic pipet. The pipet is partly filled with 1.0 M copper(II) sulfate, the tip is blocked with solidified agar–agar in 1.0 M potassium nitrate, and a copper wire has been pushed through the bulb and into the copper(II) solution.
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electrode. We find well plates particularly well-suited to minimize and organize these “other” half-cells. Use safety goggles during the experiment even though the risk of injury is very low. Copper Reference Electrode
Construction A copper reference electrode can be made as follows2: 1. Draw 1.0 M copper(II) sulfate solution into a small pipet (with a cylindrical stem) to almost fill the bulb. 2. Release a few drops of the solution and, while still squeezing the bulb, push the open end of the pipet into a solidified agar–agar salt (KNO3) gel. Release the pressure on the bulb; agar is now blocking the tip and simultaneously makes the salt bridge. Twist the pipet while pulling it out of the agar gel. (The agar in the tip must be in contact with the Cu2+ solution.) 3. Push a copper wire through the bulb and down into the Cu2+ solution. It should not touch the agar–agar salt tip. The wire can also be inserted before the agar– agar salt bridge is made. To facilitate getting the wire through the bulb, a needle could be used to make a small hole or the wire can be heated before pushing it through the bulb. Take care not to touch the wire while hot.
Between each measurement rinse the tip of the copper reference electrode in distilled water and in 1 M potassium nitrate or alternatively leave it there. One well of your culture plate is convenient for this purpose. With this electrode you can measure reduction potentials relative to Cu兾Cu2+.
Comments Agar–agar salt gel is made as follows: (i) add 1 g agar to 50 mL of cold 1 M KNO3 while stirring, (ii) heat and continue to stir until dissolved, (iii) pour the liquid into an appropriate dish (e.g., Petri dish) until it is about 1-cm deep, and (iv) leave to solidify. It is convenient to make a stock of solidified agar–agar salt gel for the salt bridge before the experiment. If left in the refrigerator, it will keep for months. If no agar is available, thick corn porridge with salt added will work as a salt bridge. It is made by adding one teaspoon (about 3 g) of corn flour to 10 mL of saturated solution of sodium sulfate, stirring, and then heating until it solidifies. It should be used within a day as it shrinks. If no copper wire is available, you may use a copper ribbon, but then you can only use the stem of a pipet (without the bulb). The disadvantage is that the electrode must be held upright to avoid spilling solution. The advantage is that this shape of electrode opens for the use of metals that are not purchased in the form of wire, for example, zinc. “Standard” Hydrogen Electrode
Construction To obtain reduction potential values comparable to the standard reduction potentials tabulated in chemistry books, you need a standard hydrogen electrode for calibration. A “standard” hydrogen electrode can be made as follows:
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Figure 2. The voltage between a hydrogen, (–)-sign, and a chlorine electrode, (+)-sign. The two mini plastic pipets with cut off stems are filled with 1.0 M HCl and placed with a snug fit in a snap cap vial containing 1.0 M HCl. Pt wires have been pushed through the top of the bulbs and act as electrode material.
1. Cut off most of the stem of two small pipets. Verify that they fit snugly together, with their stems down, into a small vial (e.g., VWR International 110622-24) as shown in Figure 2. 2. Remove the pipets from the vial and push Pt wires through the top of the bulbs and half way down the stem. Leave enough wire above the bulbs to act as terminals for battery connection. 3. Fill both pipets with 1.0 M hydrochloric acid using a capillary pipet. (If necessary, make a capillary pipet by pulling the stem of a wide stem pipet until it becomes very thin). 4. Add 1.0 M HCl to the small vial and replace the two pipets as described in 1. 5. Connect the Pt wires to a 9-V battery and electrolyze until some gas is formed in both bulbs. Assign the bulbs with + and − signs corresponding to the poles of the battery. A hydrogen electrode is in the (−)-bulb and a chlorine electrode is in the (+)-bulb. The voltage measured over this cell gives a standard reduction potential of chlorine directly. We obtained +1.36 V with an inexpensive digital multimeter (Figure 2). Electrolyzing hydrochloric acid so that the gases produced are collected in the pipets reduces the exposure to the chlorine gas.
Comments The voltage measured over the “standard” hydrogen electrode, assigned the (−)-bulb, and the reference copper elec-
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initially appears to be a 1:1 ratio of the gas volumes in the two pipets is soon (5–15 minutes) in different proportion (eventually about 1:5) as most of the chlorine gas produced is soon dissolved in or has reacted with water. Furthermore the student might also observe that the transparent plastic material becomes opaque as the chlorine dissolves in this nonpolar material, inviting further questions from interested students. A frequent question from students, as well as teachers (us included), is whether oxygen or chlorine is produced at the positive electrode, especially if standard reduction potentials are consulted. Mass spectroscopy revealed that only chlorine gas was produced during electrolysis at 9 V, proving advantageous kinetics for this reaction. Notes 1. Carbon rods can be used although they are less reliable. 2. Electrodes of other metals, for example, zinc can be made in a similar manner. Figure 3. The voltage between copper and the standard hydrogen electrode. The copper reference electrode is placed next to the hydrogen and chlorine electrodes in the solution in the vial, and the voltage between hydrogen, (–)-sign, and copper electrode is measured. (The red lead is a “left-over” from the electrolysis of HCl, as it is best to leave the system undisturbed).
trode gives a standard reduction potential of copper(II). We obtained +0.34 V with this method (Figure 3). This value is then added to the potentials measured between the copper reference electrode and the various metals in their salt solutions in order to obtain standard reduction potentials. The hydrogen and chlorine electrodes can be reactivated by repeated electrolysis if they start to fail. The attentive student will observe, if he or she continues to electrolyze hydrochloric acid for a while, that what
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Literature Cited 1. Biegler, T.; Woods, R. J. Chem. Educ. 1973, 50, 604–605. 2. Dillard, C. R.; Kammeyer, P. H. J. Chem. Educ. 1963, 40, 363–365. 3. Craig, N. C.; Ackermann, M. N.; Renfrow, W. B.; J. Chem. Educ. 1989, 66, 85–86. 4. Grønneberg, T. Elektrokjemi og Metaller; Universitet i Oslo: Oslo, Norway, 1985; p 45. 5. Damerell, V. R. J. Chem. Educ. 1930, 7, 1664–1667 6. Herron, F. Y. J. Chem. Educ. 1957, 34, A11. 7. Reimann, A. Chemie in unsere Zeit 1989, 23, 100–101. 8. Sanger, M. J.; Greenbowe, T. J. J. Res. Sci. Tech. 1997, 34, 377–398. 9. Sanger, M. J.; Greenbowe, T. J. J. Chem. Educ. 1997, 74, 819– 823. 10. Özkaya, A. R. J. Chem. Educ. 2002, 79, 735–738. 11. Birss, V. I.; Truax, D. R. J. Chem. Educ. 1990, 67, 403–408.
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