Galvanic Cells (Demonstration)

Voltaic cells    :    Electrochemistry

Illustrate how a galvanic cell works, by showing students the “chemistry version”, then an application- the clock!

Safety goggles should always be worn in the lab

.1 M ZnS04
.1 M CuS04
Leaves and dirt
Zinc electrode
Copper electrode
Salt bridge (or porous cup)
“Clock” apparatus

Set up the regular galvanic cell by placing the zinc sulfate solution in one container and the copper sulfate in another.  Insert the copper electrode into the copper solution, and the zinc in the zinc solution.  The salt bridge should connect the two.  Hook the red (cathode) clamp, of the voltmeter, to the copper and the black (anode) clamp to the zinc.  The voltmeter should read approximately 1.1V.  Then show the students an application, of use of the flowing electrons, with the “clock” apparatus, using something containing electrolytes to allow the flow of the ions, like pickles, oranges, or lots of other things.  You can show how this “shouldn’t” work by using distilled water.

The zinc is oxidized into Zn2+ ions and 2 electrons.  The voltmeter cables provide an external network for the electrons to flow to the copper side and cause reduction of the Cu2+ ions to Cu (s).  The salt bridge is what allows the flow of the ions, this is how the electrolyte works.  Since this reaction is occurring without the input of energy it is a spontaneous oxidation reduction reaction, and therefore the cell is termed “galvanic.”  The electromotive force between the two electrodes of different voltage is measured with the voltmeter.  The 2 reactions taking place at each electrode are:

Zn electrode (anode):  Zn(s)   Zn2+ (aq) + 2e-
Cu electrode (cathode):  Cu2+(aq)  +  2e-    Cu(s)

Solutions can be saved and used again.  The zinc electrode may need to be cleaned with steel wool.

Ginger Chateauneuf, 2000.