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Redox Main
Half-Reactions Balancing Redox Reactions in Neutral, Acidic, and Basic Solutions |
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Electrochemical Cells
In the previous lesson, we gave some hypothetical examples of reactions: some in which one reactant "wants" to be oxidized, and the other wants to reduced; and some in which the reactants compete to get (or compete not to get) the electrons. However, these reactions can be more than hypothetical; we can actually see them occurring by setting up an electrochemical cell, an apparatus in which oxidation and reduction take place. These devices are also called galvanic or voltaic cells. In these cells, chemical change is harnessed to generate electricity. To illustrate a galvanic cell, let's go back to our previous example:
Zn2+(aq) + Mn(s)
Zn(s) + Mn2+(aq) E° = 0.417
Notice the reactants--we need aqueous dissolved zinc ions and solid manganese. The results are solid zinc and dissolved manganese, which implies there must be a solvent for the manganese to dissolve into. Therefore, we need a water solution on both sides, a solid for zinc to plate onto, and solid manganese to dissolve. The solids in this example are called electrodes, and they may be inert (non-reacting) or may participate in the redox reaction. We also need a path for the electrons to travel through from the oxidizing manganese to the reducing zinc. Here's what a chemical cell looks like and how it works:
When the cell starts, solid manganese begins to oxidize, releasing two electrons per atom and dissolving as aqueous ions. The electrons travel across the wire, through the voltmeter (or whatever devices are connected to the cell), and down into the solid zinc electrode. There, they flow into a zinc ion, causing it to plate onto the solid electrode. Since the solution in the zinc compartment becomes less positive, manganese ions flow through the "salt bridge" to replace them. Any positive ion will do, but if the salt bridge is removed, the cell will not operate.
In our cell, electrons flow from the manganese compartment to the zinc compartment. The manganese compartment is called the anode, and is labeled using a "-" sign, because the electron flow originates here. The zinc compartment, labeled "+" because the electrons are consumed here, is the cathode. An easy way to remember this is that the "A" in anode comes before the "C" in cathode, just as the "O" in oxidation comes before the "R" in reduction--the first two letters go together ("A-O"), as do the second ("C-R").
As the reaction continues, the supply of solid manganese and ionic zinc begins to drop; the voltage also starts to decline, as the cell is not at standard conditions anymore. Eventually, the reaction will stop when the reaction reaches equilibrium; if the equilibrium constant is large, the supplies of each reactant and the accumulation of manganese ions will be the limiting factors.
However, using Le Chatlier's principle, we can lengthen the life of the cell. If we introduce a reagent to precipitate the manganese ions as soon as they come off the solid electrode, Le Chatlier's principle predicts that the reaction will try to replace the missing ions, raising the voltage as long as the supply of zinc ions and precipitating reagent lasts. However, there is nothing we can do about the lack of zinc ion or manganese metal except introduce more reactants into the cell. Eliminating the solid zinc as it forms is pointless, since the concentration of solids does not affect the reaction. Finally, heat changes could be advantageous if the reaction is exothermic or endothermic.
Modern dry-cell batteries use some of these techniques to prolong battery life. In many batteries, zinc is oxidized and ammonium ions are reduced, according to the following half-reactions:
Oxidation: Zn(s)
Reduction: 2 NH4+ + 2 e-However, ammonia and especially hydrogen are gases--if they were allowed to accumulate, pressure would build up and the battery would eventually burst! Therefore, other reactions are introduced to dispose of the gaseous products:
Consume hydrogen: 2 MnO2(s) + H2(g)
Consume ammonia: Zn2+(aq) + 2 NH3(g) + 2 Cl-(aq)Therefore, batteries include ammonium chloride (for the reduction of ammonium and the consumption of ammonia), manganese(IV) oxide (to consume ammonia), zinc(II) chloride (to consume ammonia), and a solid zinc anode to be oxidized. Adding up all these reactions gives the final redox reaction, which has an E° of 1.5 volts:
2 MnO2(s) + 2 NH4Cl(s) + Zn(s)
In the now-common alkaline battery, zinc is reduced with hydroxide to form zinc(II) oxide and water. Manganese(IV) oxide is reduced with water to form manganese(III) oxide and hydroxide; the water and hydroxide cancel in the net reaction, leaving no liquid or gaseous products. In addition, the alkaline battery has a longer shelf life than a dry-cell. Finally, it delivers an even 1.54 volts under all loads, while the dry-cell's voltage can decline under heavy load. All batteries that consume the reactants irreversibly are called primary batteries.
Also common are lead car batteries, in which solid lead, lead(IV) oxide, and sulfuric acid react to form lead(II) sulfate and water, with a potential of 2 volts. Most batteries include 6 cells, for a total voltage of about 12 volts. This reaction can be reversed, allowing the car's alternator or a battery charger to recharge the battery. Although lead batteries are large and toxic, no better alternative has been found. Such rechargeable batteries are called secondary or storage batteries.
Also important are fuel cells, in which electricity is extracted from reactants, which are continuously introduced into the cell. The best-known type of fuel cell involves the reaction of hydrogen and oxygen gas, which would normally explode. In particular, the space program has benefited immensely from this technology, as the cells are smaller and more efficient than batteries. In a fuel cell, however, the reaction delivers .9 volts per cell at moderately high temperatures (above 70 °C). KOH is used as a catalyst:
Oxidation: 2 H2(g) + 4 OH-
Reduction: O2(g) + 2 H20(l) + 4 e-Overall: 2 H2(g) + O2(g)
Not only does this reaction supply necessary electrical power for vehicles such as the Space Shuttle, but it also provides pure water used for drinking, which otherwise would have to be carried as well. Auto manufacturers are currently investigating the use of hydrogen fuel cells in cars, but they still do not deliver the power of a gasoline engine, and the rather high operating temperature presents problems. However, the development of a "clean car" would be a major breakthrough and a perfect symbol of the increasing role that chemistry plays in our lives.
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