In the Laboratory edited by
Cost-Effective Teacher
Harold H. Harris University of Missouri—St. Louis St. Louis, MO 63121
A Simple Hydrogen Electrode Per-Odd Eggen Department of Chemistry, Norwegian University of Science and Technology, 7491 Trondheim, Norway;
[email protected] The construction of an inexpensive, robust, and simple hydrogen electrode, as well as the use of this electrode to measure “standard” potentials, is described. Calculating and measuring electromotive force, EMF, of galvanic cells is central to most general chemistry curricula. To calculate cell potentials, students use a standard reduction reference scale. This scale is based on the standard hydrogen electrode, a hypothetical electrode with exact concentrations and ideal behavior (1–3). The normal hydrogen electrode is described in many textbooks and articles, but is both inconvenient and expensive (4). Instead of measuring standard reduction potentials directly, secondary reference electrodes such as calomel or silver–silver chloride electrodes are commonly used. Various secondary electrodes can be found in the literature, (5–7). Hydrogen electrodes are also reported, (8, 9), but to my knowledge, no “home-made” hydrogen electrode for direct measurements of potentials has been reported previously. Hydrogen often acts as a reducing agent, thus the electrode described here is closed to avoid hydrogen gas reacting with the substances that are to be measured. Misconceptions are abundant in electrochemistry (10), and conducting experiments may counter some of these (11), although experiments and laboratory instructions may be counter-effective due to an overload of information (12). Tables of reduction potentials refer to the standard hydrogen electrode; however, students never use this electrode in experiments, but instead some secondary reference electrode (13). This may increase the problem of understanding electrode potentials. In the experiment described here the student can directly measure the reduction potentials of many metal–metal ion pairs, as well as of some pairs of non-metallic compounds. By letting students construct a hydrogen electrode and measure potentials of various half cells relative to the (standard) hydrogen electrode, a better foundation is made before proceeding to investigate other half cell pairs. Platinum is used as electrode material in the hydrogen electrode because it is both fairly inert and because it adsorbs hydrogen readily by the application of sufficiently negative potentials (14). While platinum is the preferred electrode material, copper wire and some heat-resistant nickel wires also give satisfactory results in these trials. Silver, graphite, and glassy carbon were tested as electrode materials, but were not found suitable. In the described experiment, an inexpensive multimeter was used to measure the potentials. Such multimeters allow small currents to flow, thereby decreasing the accuracy of the experiment, but for the intended use in high school and lower-level courses at colleges or universities, the accuracy is satisfactory. Measurements with this electrode normally deviate less than 0.03 V from the standard reduction potential given in tables.
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Construction of the Electrode The tip of a wide-stem polyethene pipet is cut off and a Pt wire (e.g., 10 cm and 0.5 mm diameter) is pushed through the top of the bulb and inserted into the bulb, but not into the stem. If necessary, the hole can be made with a needle first. Enough wire is left above the bulb to act as a terminal for connection. The excess wire is curled inside the bulb (Figure 1A). The pipet is turned upside down and filled completely with 1.0 M HCl by using a capillary pipet (Figure 1B). A few drops of solution are released and, while still squeezing the bulb, the open end of the pipet is pushed into solidified agar–salt solution.1 The pressure on the bulb is released. Agar is now blocking the tip and simultaneously makes the salt bridge (Figure 1C). The electrode is now ready for use, although it has to be connected to a battery during the experiment to act as a hydrogen electrode. This will be described in the next section, together with the application for standard reduction measurements.
A
B
C
Pt wire
1 M HCl
agar–salt gel in a petri dish
Figure 1. To make the hydrogen electrode, insert a Pt wire into the pipet and fill completely with 1.0 M HCl, then block the tip with agar–salt gel.
Journal of Chemical Education • Vol. 86 No. 3 March 2009 • www.JCE.DivCHED.org • © Division of Chemical Education
In the Laboratory A
X.XX ź
Pt wire 1 M HCl
voltmeter á
copper wire 1 M solution of a copper salt
Figure 2. Measuring the standard potential to Cu/Cu2+. The inserted picture shows bubbles of H2 on the surface of the Pt wire.
á
ź
9V gel with salt
Measuring Standard Potentials In this procedure a Cu∙Cu2+ cell is used as an example (Figure 2). The experiment can be conducted with other half cells, such as Zn∙Zn2+ using the same procedure. A small test tube is filled half full with a 1 M copper salt solution (e.g., 1 M copper sulfate) and a clean wire or ribbon of copper is inserted into the solution. Enough wire or ribbon is left above the test tube to act as a terminal for the electrode. The hydrogen electrode is inserted into the test tube, the negative terminal of the voltmeter is connected to the hydrogen electrode and the positive terminal to the copper electrode (Figure 3A). Before the cell potential can be measured, hydrogen gas must be developed inside the hydrogen electrode. This is obtained by a short electrolysis as follows: (Figure 3B) The negative battery pole is connected to the hydrogen electrode and then the positive pole is connected to the copper electrode for about one second. The battery is disconnected and the “standard” potential is measured. The hydrogen electrode should not be disturbed, therefore connect the battery to the voltmeter terminals instead of directly to the electrodes (Figures 2 and 3). The tip of the hydrogen electrode is rinsed in distilled water between each measurement or alternatively, the tip is rinsed and then the agar tip is changed by squeezing out the used plug before replacing it with a new one.
B
9.00 ź
Pt wire 1 M HCl
voltmeter á
copper wire 1 M solution of a copper salt
á
ź
9V gel with salt
Hazards Hydrochloric acid is corrosive, and standard safety precautions, including the use of safety goggles, should be followed during this experiment. Cupper salts (chlorides and sulfates) are damaging to health and may cause irritation of mucous membranes. Acute poisoning may also occur, although swallowing of poisonous quantities normally leads to vomiting.
Figure 3. The voltmeter should be connected to the electrodes as shown in (A). By connecting the battery as shown in (B) for one second, hydrogen gas develops on the surface of the platinum wire. When the battery is disconnected again, the intended potential value is measured as shown in Figure 2.
Comments Before the battery has been connected to the system, the measurement will not give the intended result. When
the battery is connected, the electrolysis produces bubbles of hydrogen gas onto the platinum wire, and the electrode
© Division of Chemical Education • www.JCE.DivCHED.org • Vol. 86 No. 3 March 2009 • Journal of Chemical Education
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In the Laboratory
potential measured will change to the expected value (Figure 2). This potential will not be constant, but slowly drifting. By reconnecting the battery, for example, every minute, the potential will be fairly constant. This renewal of hydrogen gas simulates the bubbling of gas in the original normal hydrogen electrode. The battery should be connected in a way that allows the electrode to be undisturbed during the measurement. In the experiment a 9-V battery is used as an example, although a battery with less voltage may be preferable to avoid a too high overvoltage. Nevertheless, in this experience, there is no problem associated with the use of a 9-V battery. If the battery is attached with the positive pole to the Pt wire, chlorine gas will develop on the Pt surface instead of hydrogen gas. The measured potentials will then be relative to Cl2 ∙Cl– instead of H+∙H2. Platinum is expensive. A very cheap alternative is to replace it with copper wire as electrode material. However, the measured potential values will be less stable than with a platinum wire. Acknowledgment I would like to thank Lise Kvittingen, Norwegian University of Science and Technology, for both scientific and linguistic advice.
1. Smolinski, A.; Moore, C.; Jaselskis, B. The Choice of the Hydrogen Electrode as the Base for the Electromotive Series. In Electrochemistry, Past and Present; Stock, J., Orna, M., Eds.; ACS Symposium Series 390; American Chemical Society: Washington DC, 1989; pp 127–141. 2. Ramette, R. W. J. Chem. Educ. 1987, 64, 885. 3. Biegler, T.; Woods, R. J. Chem. Educ. 1973, 50, 604–605. 4. Hildebrand, J. H. J. Am. Chem. Soc. 1913, 35, 847–871. 5. Damerell, V. R. J. Chem. Educ. 1930, 7, 1664–1667. 6. Herron, F. Y. J. Chem. Educ. 1957, 34, A11. 7. East, G.; del Valle, M. A. J. Chem. Educ. 2000, 77, 97. 8. Reimann, A. Chemie in unsere Zeit 1989, 23, 100–101. 9. Eggen, P.; Kvittingen, L.; Grønneberg, T. J. Chem. Educ. 2007, 84, 671–673. 10. Schmidt, H.; Marohn, A.; Harrison, A. J. Res. Sci. Teaching 2006, 44, 258–283. 11. Sanger, M. J.; Greenbowe, T. J. J. Chem. Educ. 1997, 74, 819–823. 12. Johnstone, A. H. J. Chem. Educ. 1997, 74, 262–268. 13. Runo, J. J. Chem. Educ. 1993, 70, 708–713. 14. Rodriguez, J. J.; Herrera Melián, A.; Pérez Peña, J. J. Chem. Educ. 2000, 77, 1195–1197.
Supporting JCE Online Material
http://www.jce.divched.org/Journal/Issues/2009/Mar/abs352.html
Note 1. Agar–salt gel is made as follows: (i) add 1 g agar to 50 mL 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 ½–1 cm deep, and (iv) leave to solidify.
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Literature Cited
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Journal of Chemical Education • Vol. 86 No. 3 March 2009 • www.JCE.DivCHED.org • © Division of Chemical Education
