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.
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:
| .<-- 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 SO42-
ions to travel, maintaining equilibrium. This pathway is porous enough to allow these ions
movement, while blocking the Zinc and Copper ions.
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
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.
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.
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