In the Laboratory
The Lead–Lead Oxide Secondary Cell as a Teaching Resource Michael J. Smith,* António M. Fonseca, and M. Manuela Silva Departamento de Química, Universidade do Minho, Braga, Portugal; *
[email protected] The lead–lead oxide or lead–acid cell is indisputably one of the most successful cells ever developed. Presented in its original form by Gaston Planté in 1859, this cell evolved to become the most important commercial secondary or rechargeable cell for almost a century (1–4). Global lead–lead oxide market consumption (1999) for automotive and leisure applications, industrial stationary and vehicle propulsion (traction) batteries, and small sealed cells in electronic applications was 310 million, 48 million, and 60 million units, respectively. This gives some indication of the high commercial importance of these cells. As an example of a secondary cell the lead–lead oxide cell is generally dealt with in introductory chemistry or physical chemistry textbooks (5–8) in a rather superficial manner, with a brief description of cell or battery structure and a summary of the electrode reactions. Further details of cell structure, design developments, a more complete description of reactions, performance improvements, and models that predict the variation of cell voltage as a function of state of discharge are available in more specialized texts (2–4, 9, 10). Students frequently have difficulties with electrochemical concepts and it is often helpful to stimulate interest by relating learning objectives to familiar objects with obvious practical applications (11–13). The secondary cell, in its various chemical formulations, is present in a remarkable number of applications, making an enormous impact on daily life in modern society. In this article we describe a readily constructed laboratory version of the Planté cell that may be used as the basis of a teacherorientated upper-secondary school project or to illustrate practical aspects of secondary cell performance to physical chemistry students. In this respect the robust lead–lead oxide cell provides a hands-on contact that is not possible with any of the more advanced examples of commercial secondary cells. The time required for completing cell assembly, electrode formation, and discharge using a range of constant-load or constant current conditions is about 2–3 hours. Operation of the Lead–Lead Oxide Cell Early versions of the Planté cell were not practical in the sense that they required a prolonged “conditioning” procedure before developing a suitable surface layer of active material and a reliable discharge performance. This electrochemical conditioning consisted of a series of charges and discharges and became known as the “formation process”. Camille Fauré demonstrated in 1881 that the capacity of a new cell could be increased and stabilized more rapidly by coating lead plates with a paste of lead dioxide. Using the procedure proposed in this article conditioning is carried out at the beginning of the laboratory session by subjecting the cell to an initial charge at 20 mA during a fixed time interval and discharging to a potential of 1.6 V. This process should be repeated once or twice to form the electrodes, a process that may occupy the first 30–40 minutes of the lab session. The importance of service history on cell performance has become evident with high performance cells and quite sophisti-
negative plate clamp
positive plate clamp
O-ring seal lead electrode lead oxide electrode
sulfuric acid electrolyte
Figure 1. Schematic representation of the cell assembly.
Figure 2. Photograph of the experimental cell and bench stand.
cated charge controllers have been developed to maintain high levels of cycle efficiency under extremely demanding conditions (14). In the present application the cell voltage considered to mark the end of useful discharge service is chosen as 1.6 V and is designated as the “cut-off ” potential. Recharge should be terminated at the end of the cycle when the cell potential reaches 2.9 V. Experimental Procedure A suitable volume of 3 M sulfuric acid was transferred to a beaker, and two clean lead sheets (25 × 70 mm, 0.5 mm thickness), supported by steel clamps on a Perspex top-plate, were immersed in the acid. The top-plate, equipped with a close-fitting rubber O-ring seal (as shown in Figure 1), avoids significant loss of electrolyte should the cell be accidentally overturned. As an additional precaution, the beaker was located in a bench support that provides additional mechanical stability (Figure 2).
© 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 A
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Figure 3. Cell potential versus time during (A) a charge cycle and (B) a discharge cycle.
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Figure 4. Discharge curves obtained during constant-current discharge at (A) 10, (B) 20, and (C) 40 mA.
Figure 5. Discharge curves obtained during constant-load discharge at (A) 195, (B) 100, and (C) 44.7 Ω resistors.
Hazards
at a cut-off potential of 1.6 V and the duration of the experiment was typically 2 minutes with the test cell. The charge and discharge cycles were repeated once or twice and the cell capacity was found to stabilize to a reproducible value. During the first charge and discharge cycles the variation of cell potential with time was recorded in a table. In order for consistent cell capacity to be established after so few cycles particular care must be taken during initial preparation of the electrode surfaces. The data from the table of results were plotted to obtain a discharge curve of the cell under controlled experimental conditions. Results are shown in Figure 3 for typical charge and discharge cycles. The variation of potential during discharge is typical of primary or secondary cells and is influenced by the nature of the cell components, cell architecture and the chemical compositions of components of the test cell. The results recorded for the same cell, charged under reference conditions and subjected to three different discharges at 10, 20, and 40 mA, are shown in Figure 4. The energy converted by the cell during discharge is evaluated by calculating the area under the curve of a graph of instantaneous power developed by the cell as a function of time. The majority of real-life applications of cells involve constant-load rather than constant-current discharge and therefore the use of suitable resistors places the cell in a more representative situation (Figure 5). The constant-current source is still used to charge the cell to a convenient reference condition
Lead and lead oxide are classed as toxic materials and sulfuric acid is corrosive and may cause chemical burns and damage clothing (15, 16). Electrodes used during one lab class may be recycled for use by other students simply by subjecting them to mechanical cleaning with a steel wire brush or steel wool. This cleaning operation should be carried out within a fume hood. This allows both the cost of the experiment and the amount of material waste to be reduced to within acceptable limits. During the final stage of the cell charging procedure minute bubbles of hydrogen gas are formed on the electrode surfaces. Although the volume of gas liberated is small, hydrogen gas is highly flammable and therefore care should be taken to avoid sparks or open flames in the area. Discussion Preliminary experiments were carried out to optimize the choice of electrolyte concentration and electrode dimensions. The current intensity and duration of the charging procedure were also chosen to provide suitable cell performance. During the initial charging cycle the positive terminal plate acquired a dark brown coloration due to the build-up of a PbO2 coating and the negative terminal plate became grey due to metallic lead deposition. The first 20 mA discharge of the cell was terminated 358
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In the Laboratory Table 1. Electrical Energy Transformed by the Cell as a Function of Discharge Conditions Discharge Current/mA
Energy/J
Efficiency (%)
10
20.1
48.0
20
19.8
47.4
40
16.9
40.4
195
22.9
54.8
100
21.1
50.5
18.6
44.5
Load Resistor/Ω
44.7
between discharge experiments. The cell potential at a known time during discharge is used with the value of the load resistor and Ohm’s law to calculate the instantaneous current and power being developed by the cell. The area under the graph of instantaneous power as a function of time gives the total energy converted during the discharge cycle. Students easily arrive at the conclusion that the more rapidly electrical energy is extracted from the cell, the less efficient the energy conversion process. This conclusion is confirmed by the data included in Table 1. The results are reproducible between successive discharges of the same cell and between different cells assembled by different student-pairs over a two-semester period. The same electrode dimensions, electrolyte concentration, charge–discharge currents and resistors were used in all trials and the error of the results is estimated at about 3%. The energy efficiency of a cell charge/discharge cycle is a measure of the fraction of the energy used to charge the cell that is recovered during discharge. The energy efficiency of a secondary cell is defined by eq 1, where Edis and Ech are the cell voltages during discharge and charge and idis and ich are the currents that are applied during the discharge and charge processes, respectively. The upper limits of integration in eq 1 correspond to the times at which the charge and discharge cycles terminate.
Energy efficienccy
Edis idis dt
(1)
Ech ich dt The energy efficiency is one of the most important parameters considered by original equipment manufacturers in selecting power sources; however, a practical secondary cell or battery must also satisfy many other operational criteria. It should, for example, sustain several hundred charge–discharge cycles during its working lifetime, show high power and energy density, good charge retention, and preferably be low-cost and use only nontoxic components. Safe operation, particularly under abuse is a critical aspect. All these operational parameters are determined by electrode surface area, electrode preparation, choice of active materials, active material morphology, and cell architecture. Clearly there are many topics that students might explore as follow-up tasks to the experimental class. Conclusions The cell structure proposed in this activity provides a robust and economical basis for experimentation and may be used by different students in successive laboratories and adapted to
students of diverse academic experience. At the simplest level students may assemble a working lab-scale model of a commercially important secondary cell. As the cell described is of single compartment and single electrolyte design, assembly is simple and an operational cell can be built without recourse to specific glassware or salt bridges. The anodic and cathodic half-cell reactions are easily identified and provide real-world examples of redox reactions. By adhering closely to the cell design and experimental procedure described students can observe the expected increase in cell capacity that occurs as a result of the build-up of the active electrode material as the electrode formation process takes place. Using a simple experimental procedure students can evaluate the energy content of a low-capacity lead–lead oxide secondary cell and study the influence of discharge conditions on energy conversion efficiency. Based on the results obtained with students during laboratories the authors believe that experimentation with these secondary cells provides additional motivation and facilitates interpretation and understanding of the electrochemical basis of the processes involved. Literature Cited 1. Skundin, A. M.; Clebou, O.; Kisin, V. Chemical Power Sources; Academic Press Inc.: London, 1980; pp 198–220. 2. Handbook of Batteries, 2nd ed.; Linden, D., Ed.; McGraw-Hill: New York, 1995; pp 2.1–2.19, 24.1–25.39. 3. Vincent, C. A.; Scrosati, B. Modern Batteries: An Introduction to Electrochemical Power Sources, 2nd ed.; Arnold: London, 1997; pp 39–55, 142–161. 4. Dell, R. M.; Rand, D. A. J. Understanding Batteries; Royal Society of Chemistry: London, 2001; pp 4–5, 100–125. 5. Chemistry: Collected Experiments; Nuffield Foundation, Eds.; Longmans, Green & Co.: London, 1967; p 346. 6. Castellan, G. W. Physical Chemistry, 3rd ed.; Benjamin Cummings: Menlo Park, CA, 1983; p 398. 7. Hamann, C. H.; Hamnett, A.; Vielstich, W. Electrochemistry; Wiley-VCH: New York, 1998. 8. Atkins, P.; Jones, L. Chemistry: Molecules, Matter and Change; W. H. Freeman and Co.: New York, 1997; p 649. 9. Garche, J. Phys. Chem. Chem. Phys. 2001, 3, 356–367. 10. Treptow, R. S. J. Chem. Educ. 2002, 79, 334–338. 11. Smith, M. J.; Vincent, C. A. J. Chem. Educ. 1989, 66, 683–687. 12. Smith, M. J.; Vincent, C. A. J. Chem. Educ. 2001, 78, 519–521. 13. Smith, M. J.; Vincent, C. A. J. Chem. Educ. 2002, 79, 851–853. 14. Buchmann, I. Batteries in a Portable World; Cadex Electronics Inc: Richmond, Canada, 2001. 15. The Sigma Aldrich Library of Safety Data, 2nd ed.; Lange, R., Ed.; Sigma-Aldrich Corporation: Milwaukee, WI, 1988. 16. Toxicological Profile for Lead. Agency for Toxic Substances and Disease Registry, U.S. Department of Health and Human Services: Atlanta, GA, http://www.atsdr.cdc.gov/toxprofiles/tp13. html (accessed Sep 2008).
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