Chemistry 1100
Build your own battery
Written by L. Phillip Silverman, University of Missouri – Columbia
Modified for Chemistry 1100 by S. W. Keller
Electrolytic cells, or batteries, are a classic example of putting chemical reactions to use
in the everyday world. We have learned that many chemical reactions can be used to
generate heat, (combustion of hydrocarbons for example). In an electrochemical the
reactions are used to directly produce electricity. Many of the devices that you use
everyday, like your car cell phones, watches, pagers, are powered completely or in part
by electrochemical reactions. Today we are going to delve into what makes a battery
work, and then construct one and explore some of the variables associated with them.
Background on Electrochemical Theory
An electrochemical (or galvanic) cell, also known as a battery, is a device that produces
an electric current as the result of an electron transfer reaction. Such electron transfer
reactions are also known as oxidation-reduction, or redox, reactions. Electron transfer
occurs as one substance is oxidized, or loses electrons, while another substance is
reduced, or gains electrons. For example, if a piece of zinc metal were immersed in a
solution containing copper(II) ions, the zinc would spontaneously lose electrons while
the Cu(II) would spontaneously gain electrons. This process can be expressed as two
half-reactions that sum to yield the overall reaction:
Zn (s) → Zn+2 (aq) + 2 e- (oxidation)
Cu+2 (aq) + 2 e- → Cu (s) (reduction)
NET: Zn (s) + Cu+2 (aq) → Cu (s) + Zn+2 (aq)
Any spontaneous redox reaction can be harnessed to produce electrical energy under
the right conditions. The problem with simply dropping a piece of zinc metal into a
solution of Cu(II) is that the electrons provided by the zinc move directly to the aqueous
Cu(II) ions without doing any work. In order to create a useful battery, the two half
reactions must be physically separated so that the electrons will flow through an
external circuit as shown in Figure 1. A salt bridge is necessary for charge balance: in
this case nitrate (or sulfate) ions flow from the copper to the zinc compartment, to
balance the flow of electrons (which is in the opposite direction).
The following shorthand notation can represent the electrochemical cell shown in
Figure 1:
Zn (s)| Zn+2(aq) || Cu+2 (aq) | Cu (s)
In this type of “line notation”, the components at the site of oxidation (the anode) are
listed on the left; at the site of reduction (the cathode), on the right; and the central
double vertical lines represent the salt bridge. A single vertical line indicates a solid (s)
is immersed in the solution (aq).
Chemistry 1100
Electrons that are generated at the anode of an electrochemical cell (the oxidation
reaction) are driven toward the cathode where they are used in the reduction reaction.
The actual voltage that is obtained by any electrochemical cell depends on a number of
variables including the temperature, the composition of the two electrodes, and the
concentrations of the solutions on each side. We will be changing the solution
concentrations and determining the effect on the voltage that the cell is able to generate.
You will be making 4 separate cells, one at a time, and measuring the voltages from
each.
Procedures:
Part I: Solutions
1. In a 100 mL beaker, dissolve 12.5 g CuSO4●5H2O in 50 mL water. This is your
1.0 M Cu2+ solution.
2. In another 100 mL beaker, take 5 mL of the 1.0 M Cu2+ solution that you just
made and add 45 mL water. This is your 0.1 M Cu2+ solution.
3. In a third 100 mL beaker, dissolve 12.9 g ZnSO4●7H2O in 50 mL water. This is
your 1.0 M Zn2+ solution.
Chemistry 1100
4. In a fourth 100 mL In another 100 mL beaker, take 5 mL of the 1.0 M Zn2+
solution that you just made and add 45 mL water. This is your 0.1 M Zn2+
solution.
Part II: Salt Bridge
1. Weigh out about 2 g of agar.
2. On a hot plate, boil 100 mL of the 0.1 M KNO3 solution that is
provided.
3. Remove the KNO3 solution from the heat and quickly add your
2 g of agar, stirring until the agar dissolves.
4. Fill a U-tube with the agar solution before it cools, leaving
about a 1/8 inch of air space at each end of the tube. Insert
cotton plugs into each end, leaving some cotton protruding
from each end. Pour some hot agar solution poured over the
cotton plugs so that the cotton is saturated. Do not remove the
cotton plugs until you are completely finished with the experiment!! Allow
your salt bridge to air cool (no water baths!). The agar will solidify upon
cooling.
Do not turn the tube upside-down until the agar has solidified completely!!
Part III: Assembly and Operation of Cell
1. Use a U-tube as a salt-bridge to link the two beakers prepared in step (1).
2. Obtain a piece of copper and a piece of zinc, and clean the surfaces with an
emery board or sandpaper.
3. Insert the zinc strip into the beaker of 1.0 M ZnSO4 and the copper strip into
the beaker of 1.0 M CuSO4. Obtain a voltmeter (you will need student ID
card to check out the voltmeter from the stockroom) and attach the positive
lead to one strip and the negative lead to the other strip using alligator
clips.
4. Read and record the voltage of the galvanic cell. You may need to reverse
connections to achieve a positive voltage, just keep track of which half-cell is the
anode and which is the cathode. Unhook your cell after taking the reading.
Record your results on the board and compare your readings with those
of your classmates.
DO NOT MOVE ANY PART OF YOUR GALVANIC CELL or your voltage will fluctuate.
5. Add 0.85 grams of silver nitrate (AgNO3) to a 200 mL beaker. Add 50 mL of
water and stir the solution until the salt completely dissolves in the water. You
have just created a 0.1 M AgNO3 solution.
6. Make the remaining cells described in the chart listed below.
Chemistry 1100
7. Make another galvanic cell this time using the 0.1M CuSO4 solution with
copper metal on one side and the 1.0 M CuSO4 solution with copper metal on
the other end of the cell.
8. After measuring the first five cells, you will make four more galvanic cells this
time using the 0.1 M AgNO3 solution with silver metal on one side other
metal/solution combinations as described in the chart below
At this point you SHOULD have the voltages for 9 cells. Do not proceed until
all 9 cells have been measured!!
9. Dispose of the solutions in the appropriate containers and dismantle the cell.
Used salt bridge material goes into the container identified as “unwanted salt
bridge materials”. The metal strips/wires can be reused.
1.0 M ZnSO4
0.1 M ZnSO4
“ALL-COPPER”
CELL
1.0 M CuSO4
1)
2)
1.0 M Cu||0.1 M Cu
With Cu Metal
electrodes on both sides
0.1 M CuSO4
3)
4)
5)
1.0 M CuSO4
0.1 M AgNO3
6)
0.1 M CuSO4
7)
1.0 M ZnSO4
8)
0.1 M ZnSO4
9)