Oxidation and Reduction

Batteries, also known as voltaic cells*, galvanic cells*, and electrochemical cells, play an important role in electronics. A battery can also refer to a collection of these cells put together. These devices produce a current at a constant voltage, which makes them very useful. The current is produced by a chemical reaction known as a redox reaction.

A redox reaction is the short name for an oxidation-reduction reaction. We think the scientists who coined the phrase thought "oxred reaction" sounded too funny to be scientific. The terminology used with redox reactions can be very confusing, so you might want to memorize this little table.

Oxidizing Agent	The molecule accepting electrons.
Reducing Agent	The molecule providing electrons.
Oxidation	A molecule provides electrons and is oxidized.
Reduction	A molecule accepts electrons and is reduced.

You can look at this table of reduction potentials to see the chemical equations for some redox reactions. Don't worry if you don't understand what the equations mean, just look for patterns and similarities between the equations. (Ignore the right hand column for now.)

Redox reactions involve the transfer of electroncs from one type of molecule to another. There are two main laws that govern this type of reaction:

1. The reducing agent is always oxidized, the oxidizing agent is always reduced
2. If a substance is oxidized, another must be reduced.

You may be thinking to yourself, "You know, I wouldn't be able to recognize a redox reaction if I saw one." Don't despair! You can always recognize a redox reaction by looking for a changing oxidation number. The oxidation number is the charge the atom appears to have when we apply a set of rules to that substance. Here are the rules:

1. A free element's oxidation number is 0.
2. If an ion only contains a single element, the charge on that ion is the oxidation number
3. The oxidation number of H is usually, but not always, +1
4. The oxidation number of O is usually, but not always, -2
5. The sum of a neutral compound's oxidation numbers must be 0
6. The sum of a polyatomic ion's oxidation numbers must equal the ion's charge.

To balance a redox reaction, first break the process into half reactions. In our earlier example we used:
Cu²⁺(aq) + 2e^- ==> Cu(s)
Zn(s) ==> Zn²⁺(aq) + 2e^-
Each half reaction should be balanced by mass and charge. As you can see, we tossed a couple of + 2e^-s to make everything balance.
Next we would multiply each reaction by a factor that would make the number of electrons given off equal the number received. (While this is not a problem in this more basic example, this is quite a common procedure.) Then we add the two half reactions to give us a fully balanced redox equation. It is slightly more difficult if the compounds are in a acid or base solution, you may want to consult a chemistry textbook if you are interested.