Clackamas Community College




Contact instructor:

Eden Francis

Physical Science
19600 Molalla Avenue
Oregon City, OR 97045
(503) 594-3352
TDD (503) 650-6649

Lesson 9: Voltaic Cells

The tendency for the electrons to flow from one chemical to another, such as from the zinc metal to the copper ion as shown here, is something that can be channeled and controlled. Channeling the flow of electrons is what we will take up here.

Voltaic cells use spontaneous electrochemical reactions. (Remember from an earlier section of thelesson that electrolytic cells use nonspontaneous, or forced, electrochemical reactions.)

sponrx1.JPG (2942 bytes)

Zn rtarrow.gif (850 bytes) Zn2+ + 2e-      0.76v
Cu2+ + 2e- rtarrow.gif (850 bytes) Cu     +0.34v
Zn + Cu2+ rtarrow.gif (850 bytes) Zn2+ + Cu    1.10v


Copper-Zinc Voltaic Cell

The apparatus shown here is used to channel electron flow in this reaction and is called a voltaic cell. The zinc metal is on the right side and it is providing electrons to the blue copper ions in the beaker to the left. However, to get there the electrons have to travel through a wire. We can see that they are doing so because the voltmeter that they must pass through shows a reading.

Copper-Zinc Voltaic Cell

This voltaic cell has copper metal and copper sulfate solution in the left hand beaker; zinc metal and zinc sulfate in the right hand beaker; probably potassium sulfate in the salt bridge. With the wires connected to the voltmeter, you can see that there is a voltage reading; electricity is flowing through this cell. It's a voltaic cell because the chemical reaction is causing the flow of electric current.

Here you can see the importance of the salt bridge. In the top picture, the salt bridge has been removed and and the voltage reading is zero showing the there is no current flowing. In the bottom picture, the salt bridge is replaced and the voltmeter again shows a reading.

Voltaic Cell with and without a salt bridge.

This same voltaic cell is diagrammed for you here (and in example 20 in your workbook). In the diagram, copper is on the right and zinc is on the left.

Diagram of Cu-Zn voltaic cell.

Zinc metal is reacting with the copper ion to form zinc ion and copper metal. The zinc metal is on one side and the copper ion in solution is on the other side in the other beaker. Consequently, the two chemicals are not in direct contact with one another.

Even so, the tendency for the zinc metal to lose (or transfer) electrons to the copper ion still exists. That transfer is made possible by connecting a wire between the zinc metal and the copper metal. The electrons go from the zinc over to the copper metal where they can react with the copper ions in solution.

Over a period of time, the zinc electrode will dissolve and increase the concentration of the zinc ion solution. The copper ion will plate onto the copper electrode and thus the concentration of the copper ion in solution will decrease. As this happens, the reaction slows down and the voltage decreases.

Zinc is the anode because that is where oxidation is occurring. The oxidation half-reaction is Zn rtarrow.gif (850 bytes) Zn2+ + 2e-. Those electrons go over to the copper side where they react with copper ion and change it into copper metal (Cu2+ + 2e- rtarrow.gif (850 bytes) Cu), which is the reduction half-reaction that takes place at the cathode.

Diagram of CuZn voltaic cell with half-reactions.

The voltage for such a cell can be calculated using the standard oxidation potential list. Working that through, as we have done before, the voltage is 1.10 volts.

Zn rtarrow.gif (850 bytes) Zn2+ + 2e-      0.76v
Cu2+ + 2e- rtarrow.gif (850 bytes) Cu     +0.34v
Zn + Cu2+ rtarrow.gif (850 bytes) Zn2+ + Cu    1.10v

However, if you measure the voltage of a cell like this using a voltmeter, you will likely not get that particular voltage. The reasons are several.


One is that quite often solutions are not kept at 25oC, particularly in this lab.


Another reason is that the concentrations of all the solutions are probably not one mole per liter.


Also, the voltage measured by a voltmeter depends on the electrical resistance of the cell. So the electrical resistance in the connections (especially bad ones), the wire, the voltmeter, and the solutions will tend to cause the voltage to be less than what you calculate it to be under standard or ideal conditions.

We have a cell like this set up in the lab. You should look at it, experiment with it, remove and replace the salt bridge, check the voltage, and so forth, when you are in the lab.


Dry Cells

If you have never cut open a dry cell battery to see what it looks like inside, this picture gives you some idea of what it looks like. On the left is the metal case with some cardboard to keep it from contact with the reactive chemicals. Usually just the top and bottom are in contact with the chemicals, not the sides of the case. The actual workings of the battery are shown on the center and right side. A zinc metal cylinder has been cut away and put in the center of this picture. The zinc case is what becomes oxidized and provides the electrons, so it is the anode. Inside of that is a pasty mix of manganese dioxide and a solution of ammonium chloride. Even though it is called a dry cell, it is moist inside.

Cut away dry cell.

The manganese dioxide is what takes the electrons from the zinc. That particular chemical is not on your SOP list. If you want to know about it, ask and we can help you fill it in. So manganese dioxide is the chemical on the right side of an SOP list that is reacting with the zinc from the left side of the oxidation potential list. If you calculate the voltage for this cell under standard conditions, it comes out to be 1.98 volts.

            Zn rtarrow.gif (850 bytes) Zn2+ + 2e-               +0.76v
Mn2+ + 2 H2O rtarrow.gif (850 bytes) MnO2 + 4 H+ + 2e-   -1.22v
               Zn rtarrow.gif (850 bytes) Zn2+ + 2e-             +0.76v
MnO2 + 4 H+ + 2e- rtarrow.gif (850 bytes) Mn2+ + 2 H2O  +1.22v
MnO2 + 4 H+ + Zn rtarrow.gif (850 bytes) Mn2+ + 2 H2O + Zn2+   +1.98v

If you are familiar with dry cells, you know that they are about 1½ volts. So the internal resistance and the variation in the concentrations cause a drop in the voltage from the theoretical standard.


Lead Acid Battery

This picture shows the electrodes from a lead acid battery.


The anode is the grey electrode on the left. It's lead and it reacts to become lead sulfate, which is a light grey-colored material.


The cathode is made of lead that is coated with lead dioxide, which is a reddish-brown chemical. As the lead dioxide reacts, taking on electrons, the lead goes from a +4 state to a +2 state. The oxygen combines with the hydrogen to form water, and the lead in the +2 state combines with the sulfate to form lead sulfate. 

The lead acid battery is kind of interesting in that the product of both of the half-reactions is the same, lead sulfate.

Lead acid battery electrodes.

The lead acid battery involves, of course, an oxidation half-reaction and a reduction half-reaction (shown below). The oxidation half-reaction is on your list at 0.36 volts. Lead metal gives up electrons and combines with sulfate ion to form lead sulfate. The other half-reaction is that lead sulfate (PbSO4) plus water forms lead dioxide (PbO2) and 4 hydrogen ions and a sulfate ion and 2 electrons. Its voltage is -1.70 volts. The standard voltage for the overall reaction is calculated to be 2.06 volts.

Pb + SO42- rtarrow.gif (850 bytes) PbSO4 + 2e-                            +0.36v
PbSO4 + 2 H2O rtarrow.gif (850 bytes) PbO2 + 4 H+ + SO42-   + 2e-    -1.70v
Pb + SO42- rtarrow.gif (850 bytes) PbSO4 + 2e-                            +0.36v
PbO2 + 4 H+ + SO42-  + 2e- rtarrow.gif (850 bytes) PbSO4 + 2 H2O     +1.70v
Pb + PbO2 + 4 H+ + 2 SO42- rtarrow.gif (850 bytes) 2 PbSO4 + H2O     +2.06v

If you are familiar with lead acid batteries, you know that they run at 6 and 12 volts. So several of these cells have to be put in sequence with one another to increase the voltage up to those amounts.


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