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For this demonstration you will require some aluminum foil to be used as an anode in your voltaic cell. You will also need a cathode. I am fortunate enough to have had mercury-silver cathodes installed in my mouth when I was a child. That way, wherever I am, I am able to construct my own batteries. Only one more thing is needed: an electrolyte. This I also keep in my mouth. If you are similarly equipped, roll up a ball of aluminum foil and pop it into your mouth. Move it from side to side, chewing it thoroughly. If the aluminum happens to connect with a cathode, you will perceive an electric shock. Congratulations! You have constructed a voltaic cell.
Such a cell is not entirely satisfactory, however. While it produces a decent voltage, the current produced is quite small. Similar proof-of-concept batteries may be constructed from coins and salt-water or lemon juice. It may even be possible to run a tiny electronic clock from one of these, as such clocks do not require much in the way of current. But for lighting light bulbs and running motors and closing relays we need a lot more current than can be provided by a battery made from coins. The key to high currents is electrode surface area.
The battery you will build is a variation on the one described in the previous section. I chose aluminum as the anode because it is commonly available and because its standard reduction potential is rather large and negative. Ordinary table salt may be used as an electrolyte, but I have found that washing soda, sodium carbonate, produces a cell with a higher current. In the previous section we discussed platinum as an inert electrode. Such an electrode allows water itself to serve as an oxidizing agent. Recalling that the oxidizing agent is, itself, reduced and that the site of reduction is called the cathode, we see that platinum served as the cathode. Platinum is expensive, however, and so we require a cheaper material for our cathode. For electrical applications, carbon often serves as "poor man's platinum." We would like a form of carbon with a large surface area so that electrons will be more likely to cross our wire bridge than to move directly from the aluminum to the electrolyte. An inexpensive, high-surface-area form of carbon is met with in activated charcoal. Activated charcoal is porous, like a sponge, and is sold for use in water filters for aquariums. Ordinary charcoal is electrically conductive, but lacks the large surface area of activated charcoal. Since it is inert, it may be re-used to make generations of voltaic cells.
You will also need a convenient container for your battery and I find that a small stainless steel mixing bowls serve well. Three mixing bowls of the same size will allow you to construct a battery of two cells in series. The larger the mixing bowls, the greater the current from your battery, but even small, 2-cup bowls will power a small motor or telegraph sounder.
Begin by preparing a saturated solution of washing soda. Fill a mixing bowl half-full of water, add some washing soda, and stir until it is dissolved. Keep adding washing soda and stirring until no more will dissolve and solid sodium carbonate settles to the bottom of the bowl. Fill each of your remaining bowls about 1/4 full of activated charcoal. Add enough washing soda solution to each bowl to make a slurry with the activated charcoal. The charcoal will foam; air is displaced from the pores as the charcoal absorbs the solution. You now have two bowls, 1/4 full of charcoal and electrolyte. We are about to add the aluminum anode, but we need a way to electrically insulate the aluminum from the charcoal while still allowing the ions in the electrolyte to move about.
To do this, line each of the two bowls with two thicknesses of paper towel. This will keep the charcoal separate from the aluminum while allowing the electrolyte to flow between the two. Add enough washing soda solution to completely wet the paper towels. Now line each bowl with aluminum foil, taking care that the foil does not tear the paper towels and that it does not touch the mixing bowl beneath. You are now ready to connect your two cells in series.
Place one of your cells inside the other, just as if you were stacking the bowls for storage. The aluminum anode of the lower cell should make contact with the steel-charcoal cathode of the upper one. Empty the third mixing bowl of leftover washing soda and nest it within the upper cell. You now have three bowls nested one inside the other. If you were to look at their contents in cross section, it would go steel-charcoal-paper-aluminum-steel-charcoal-paper-aluminum-steel. Use both hands to press the upper and lower bowls together. By compressing the charcoal, you reduce its resistance and thereby increase the current produced by your cells. Finally, use a strong rubber band to hold the bowls tightly together.
Inexpensive volt-ohm meters are now widely available. Use a meter to measure the voltage between the lower bowl and the middle bowl, the middle bowl and the upper bowl, and lower bowl and the upper bowl. Record these voltages in your notebook and compare the voltages of the individual cells to that of the two of them in series. If your meter measures milli-amps as well, record the current output of each individual cell and compare them to that of the two in series. The voltage of your battery should be about 2 V and the current should exceed 50 mA.
There is a break-in period with these cells as aluminum hydroxide builds up around the anode. Be patient before concluding that your cells are not operating correctly. If one or both of your cells fail to produce any output, it is likely that the paper is torn, shorting anode to cathode; the electrons are able to move from anode to cathode without going through your meter. If both cells perform, but one performs better than the other, it is likely that the under-performing cell does not contain enough electrolyte. Dismantle the cell and add some more sodium carbonate solution.
The real test of your battery is whether or not it is strong enough to power an electrical device. You may try lighting a small flashlight bulb, running a small motor, or activating a telegraph sounder. You may even build your own electromagnet. Simply wind 50-100 turns of insulated wire around an iron nail. The wire must be insulated; "bell wire" or "magnet wire" can be found at hobby electronics stores. Magnet wire may also be salvaged from "dead" motors. With a knife, strip the insulation from the two ends of the wire; connect one end to the lower bowl and one to the upper bowl of your battery. When connected to your battery, your electromagnet should be strong enough to lift a paper clip.
Your battery will remain at full strength for hours but not for days. Eventually the aluminum anode will be consumed and the electrolyte will become saturated with aluminum hydroxide. The charcoal, however, is inert. When your battery is dead, or when you are tired of it, throw the aluminum foil and paper towels away. Drain the electrolyte from the charcoal and wash the solution down the sink. Wash the charcoal with a few portions of fresh water, drain, and allow the charcoal to dry. The mixing bowls and charcoal may be used to make more batteries later on. I would not, however, put the used charcoal in my aquarium filter; the residual alkali and aluminum hydroxide would probably not be good for fish.
The aluminum-alkali battery is far from optimal. One of the most practical cells of all time is the Leclanché cell, developed by Georges Leclanché in 1866. Batteries up to that time had consisted of jars filled with liquid, which were subject to breakage, spillage, and leakage. Leclanché's innovation was to produce a "dry cell," which might more accurately have been called a "moist cell." The Leclanché cell consists of a carbon/manganese dioxide cathode, surrounded by a paste of ammonium chloride, separated from the zinc anode by a layer of paper. The redox reaction is:
Zn(s) + 2 MnO2(s) + 2 NH4Cl(aq)
= ZnCl2(aq) + Mn2O3(s) + 2 NH3(aq) + H2O(l)
Unlike the aluminum-alkali cell, the Leclanché cell produces no gases and so it may be hermetically sealed. As the cell ages, however, the ammonium chloride gets used up and the cell voltage drops. This drawback has been addressed by the alkaline dry cell.
The alkaline dry cell uses the same anode and cathode as the Leclanché cell and consequently has very nearly the same EMF. Rather than using ammonium chloride as the electrolyte, however, it uses potassium hydroxide. The redox reaction is:
Zn(s) + 2 MnO2(s) + H2O(l) = Zn(OH)2(s) + Mn2O3(s)
Because the electrolyte is not consumed in the reaction, its concentration remains constant as the cell ages. Consequently the EMF of the alkaline dry cell remains nearly constant as it ages.
All of the cells we have discusses so far have been primary cells; once the anode is consumed in the reaction, the cell is dead and must be discarded. A secondary cell is one which can be recharged, that is, one in which the redox reaction may be run backwards to restore the original state of the cell. We will re-meme-ber this I-dea in Chapter 25.
Record in your notebook the voltages and currents produced by your individual voltaic cells and by the two cells connected in series. Your battery passes muster when it can operate an electrical device.
In early versions of this project I used copper dishwashing pads as the cathode. The I-dea for using activated charcoal came to me from Reference .