voltaic cell

Voltaic cells work by utilizing a spontaneous redox reaction. The cell forces the reaction to take place through an external pathway rather than directly between reactants.

Because redox reactions are the transfer of electrons, when voltaic cells force the reaction to occur through an external pathway, it creates a flow of electrons from one point to another (electricity). The point from which the electrons come is the anode. Anodes are the negative electrode on the battery because electrons are negatively charged, and the anode is the part of the cell where the electron flow originates. The cathode is the positive electrode of the battery. Here, electrons react to oxidize some species.

The voltaic cell is able to produce electricity by using species which natually tend to donate and receive electrons. Thus, the species that spontaneously gives up electrons is reduced and the species that spontaneously gains electrons is oxidized, driving the reaction foward.

A voltaic cell can also be called a galvanic cell, a galvanic battery, or a voltaic battery.

Definition

What is a voltaic cell? A voltaic cell is a device in which chemical energy is converted into electric energy. This is accomplished by submerging two dissimilar metals in an electrolyte, connecting the two peices of metal with a wire or other conductive substance, and a salt bridge. The voltaic cell is also reffered to as a galvanic cell, paying homage to the two 18th century scientists who pionerred the reasearch into this topic, Alessandro Volta and Luigi Galvani.

Introduction

Voltaic cells also rely on an important type of reaction: oxidation-reduction reactions, also called redox reactions. In an oxidation-reduction reaction, two substances interact by transferring electrons. Oxidation is defined by a loss of electrons, Reduction is a gain of electrons. This can be remembered by the mnemonic "LEO the lion says GER". Loss of electrons - Oxidation; Gain of electrons - Reduction. In Voltaic cells, the reagents more often than not consist of a metal plate (like zinc or copper) in a solution that contains the ionic form of that metal (Zinc Sulfate, ZnSO₄ or Copper Sulfate). This gives us a set up like the following, rendered in full ascii-art glory:

                        Conductor
        |============================================|
        |                                            |
        | .<-- Copper(Cu)                            |.|
        , |         _____________________________    , | <-------Zinc (Zn)
\      .  |         || /      Salt Bridge    \ ||   .  |            /                
|      |  |         || |                     | ||   |  |            |
|/\/\/\|\/|/\/\/\/\/||\|                     |/||\/\|\/|/\/\/\/\/\/\|
|      |  |            |                     |      |  |            |
|      |  .            |                     |      |  .            |
|      |.'             |                     |      |.'             |
|                      |                     |                      |
|                      |                     |                      |
|        CuSO4         |                     |                      |
|                      |                     |       ZnSO4          |
|                      |                     |                      |
|                      |                     |                      |
|______________________|                     |______________________|

The last important part is the salt bridge. The salt bridge serves as a pathway for the exchange of electrons in the opposite direction of the conductor (think about it like a chemical circuit). In this case, it would be a pathway for the SO₄^2- ions to travel, maintaining equilibrium. This pathway is porous enough to allow these ions movement, while blocking the Zinc and Copper ions.

Now what?

Now that all this is set up, what is happening? Since Zinc wants to lose electrons more than Copper, the Zinc will lose it's electrons and the Copper will gain electrons to maintain the balance. Thus, the Zinc is oxidized and the Copper is reduced. Knowing this, we can also call the Zinc metal strip the anode and the Copper metal strip the cathode. A useful way to remember this is through yet another mnemonic..."Cats have paws"(pos, for positive). Since the copper is gaining electrons, it can be thought of as positive, and therefore it is the cathode. The real definition of the cathode is the place where reduction occurs, and the anode is where oxidation occurs. So how are these electrons actually transferred? Looking at each solution, we can derive two equations:

Zn (s) --> Zn²⁺(aq) + 2e^-
Cu²⁺(aq) + 2e^- --> Cu (s)

These are called the half equations of the cell, because if we put them together, we get a description of what is happening in the entire system: zinc plus aqueous copper ions reacts to form copper and aqueous zinc ions. The sulfate is ignored because it is neither oxidized nor reduced, it is simply an electrolyte, helping to transfer electrons and maintain equilibrium. True to our equations, if we set this situation up in real life, we would discover copper beginning to collect on the copper plate, and the zinc plate diminishing in size. The reaction obviously stops when all the zinc is used up, or all the copper ions have turned into solid copper.

Calculating Cell Potential

The first (and probably easiest) thing we can calculate is the cell potential, or voltage. First, we look at our trusty sheet of standard reduction potentials (everyone has one of those, right?) and find our half equations. Looking through, we find:

Cu²⁺(aq) + 2e^- --> Cu (s) E⁰ = +0.340V
Zn²⁺(aq) + 2e^- --> Zn (s) E⁰ = -0.763V

Notice how the half equation for the Copper yields a more positive potential? That means that this one will occur as written, as a reduction. This is how we knew the copper would be reduced, and the zinc oxidized. To calculate the reduction potential for the entire cell, we just add these two, but we must flip the zinc equation, because it doesn't concur with our original half equation. When we flip it, we also flip the sign of the potential, so this gives us .34 + .763 = 1.103 V.

note: this simplification assumes STP and equal concentrations of solution. For more complex voltaic cells, one will need to use the Nernst Equation. YMMV.

Uses

I think everyone can come up with the obvious use - batteries! Today, millions of devices contain batteries, and it's all thanks to small, portable voltaic cells. We have lead-acid batteries in our cars and alkaline batteries in our mp3 players. There are so many different types of redox reactions, batteries could be made out of many, many different elements. But it's finding which elements or compounds last the longest or produce the highest voltages that scientists look for.

Some other uses for voltaic cells are pH detectors and copper/silver/gold plating.

With some help from:
http://www.chem.vt.edu/chem-ed/echem/redox-std-potentials.html
http://hyperphysics.phy-astr.gsu.edu/hbase/chemical/electrochem.html