Names: _________________________________
___________________________________
Report for experiment 21G
Electrochemistry: Galvanic Cells and Batteries
Introduction to electrochemical cells (a.k.a. galvanic cells or batteries):
The energy produced by an electrochemical cell is due to a spontaneous redox reaction. An
electrochemical cell consists of two half-cells. Each half-cell consists of an electrode (usually metal)
submerged in its own salt. For example, the half-cell on the right in the figure below is made of a piece of
copper submerged in a solution of copper nitrate. An oxidation reaction (loss of electrons) occurs in one
of the half-cells and a reduction reaction (gain of electrons) occurs in the other half-cell. Each of these is
called a half reaction. The overall cell reaction is the sum of the two half reactions. The electrode at
which oxidation takes place is called the anode. The electrode at which reduction takes place is called the
cathode. The two half cells are connected in two ways. The electrodes are connected by a wire and the
salt solutions are connected by a “salt bridge”. Electrons are lost by the anode and flow through the wire
to the cathode. The electrons flow through the cathode until they touch the salt solution where they are
taken up by cations in the solution. Thus, the metallic anode is oxidized and the cations in the cathode
compartment are reduced. The cations resulting from the oxidation reaction at the anode enter the salt
solution in the anode compartment. The uncharged metal atoms resulting from the reduction reaction
stick to the cathode. Aqueous ions flow between the two half cells through the salt bridge to complete the
electrical circuit. The cell potential (also called cell voltage) is basically the same thing as G. It tells you
how much work a battery can do. It is the “driving force” for the reaction. It tells how far the reaction is
from equilibrium. A positive cell potential means that the forward reaction is spontaneous. (This is the
opposite of the sign convention for G.) When measuring the voltage of a cell with a voltmeter, the
reaction does not take place because the voltmeter pushes the electrons back toward the anode. An
ammeter reads the current (rate of current flow in Amps = coulombs/sec) as electrons flow from the anode
to the cathode.
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Part A: Activity Series and Simple Galvanic Cells
Step 2 (activity series). In the table below, indicate which combinations of metal and metal ion resulted in a
reaction (Yes for reaction; No for no reaction).
Cu2+
Ag+
Zn2+
Cu
Ag
Zn
For each reaction that occurred above, write the net ionic equation. Include states.
Based on your results, list the 3 metals in order of reactivity (i.e. activity, ease of oxidation).
Most active metal _______________ > _______________ > ________________ Least active metal
(easiest to oxidize)
(hardest to oxidize)
List the 3 metal ions in order of reactivity (ease of reduction)
Easiest to reduce _______________ > _______________ > ________________ Hardest to reduce
Notice the inverse relationship: If a metal is easy to oxidize, then its ion is hard to reduce.
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Step 3 (electrochemical cells). Set up the cell assigned to your group as described in the procedure and measure
the voltage. If the voltage is negative, switch the wires that are connected to the metal electrodes. Record the
measured potential (voltage) in the space provided and indicate which metal is connected to the red wire (this is the
cathode where reduction takes place). Examine the cells made by other groups and record their data. Draw and
label each of the cells that you built in this step. Include all of the following in your drawings:
1. Two beakers, each with a salt solution and a piece of metal. The metals should be connected via wires from
the multimeter and the solutions should be connected by a salt bridge.
2. Label each electrode in 3 ways: the formula for the metal; anode or cathode; + or -.
3. Use arrows to show the direction of electron flow through the wires and the direction of ion flow in
solution and through the salt bridge.
4. Below each beaker, write the half-reaction occurring in each half-cell. Indicate whether this is oxidation or
reduction.
Cell 1: measured potential _______________ V
Red wire is connected to ___________________
Cell 2: measured potential _______________ V
Red wire is connected to ___________________
Cell 3: measured potential _______________ V
Red wire is connected to ___________________
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Steps 4 ad 5 (concentration effects; Le Chatelier’s principle).
This part of the experiment will be done as a demonstration by your professor. In preparation for the demo,
answer all of the questions on this page.
Write the net ionic equation for the overall reaction that occurs in Zn/Zn2+║Cu/Cu2+ cell (same as one of the
equations you wrote for step 2 above):
The voltage (cell potential, Ecell) provided by an electrochemical cell is described by the Nernst equation, which is
discussed in detail in your textbook. For the Zn/Zn2+║Cu/Cu2+ cell, the Nernst equation is:
E cell  1.10 V 
0.0592 V
[Zn 2  ] (at 25oC)
log
2
[Cu 2  ]
This equation should look familiar to you. Recall that a positive cell potential (voltage) means that the reaction will
occur spontaneously in the forward direction. Like G, the magnitude of the cell potential tells you how far the
reaction is from equilibrium. (Zero volts means the cell is dead and the reaction is at equilibrium.) G or Ecell are a
measure of the “driving force” for the reaction. Like a stretched spring, a large G or Ecell means that there is a
large “force” that wants to drive the reaction towards equilibrium.
Use the Nernst equation to predict what should happen to the voltage when the following changes are made to your
cell (circle one):
a. add water to the beaker with the Zn2+ solution
b. add CuSO4(s) to the beaker with the Cu2+ solution
c. add ammonia to the beaker with the Cu2+ solution
(note: ammonia reacts with copper ion to form a complex ion)
decrease
decrease
decrease
no change
no change
no change
increase
increase
increase
Use LeChatelier’s principle to make the same predictions. If LeChatelier predicts that equilibrium would shift to
the left (toward reactants), then you would predict that the voltage would decrease (and vice versa).
a. LeChatelier predicts that the equilibrium would shift to the ____________ so the voltage should ____________
b. LeChatelier predicts that the equilibrium would shift to the ____________ so the voltage should ____________
c. LeChatelier predicts that the equilibrium would shift to the ____________ so the voltage should ____________
Note: In a freshly made electrochemical cell, the reaction is far from equilibrium. When you measure voltage,
the reaction does not actually occur so, when you make the changes above, the reaction does not start out at
equilibrium nor does it shift. However, the voltage tells you whether you are closer to or further from
equilibrium after a change is made. It tells how much the reaction “wants” to proceed. (When reading voltage,
the electrons in the anode are like a racehorse in a starting gate!)
Your professor will now make the changes and you can see if your predictions are correct!
Original cell ____________ V After diluting the Zn2+ solution ____________ V
Original cell ____________ V After adding solid CuSO4 to the Cu2+ solution ___________ V
Original cell ____________ V After adding ammonia to the Cu2+ solution ____________ V
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Write the balanced half-reactions and overall reactions for each of the cells that you built in step 3. Use the table of
standard reduction potentials at the end of this document to predict the standard cell potentials.
Eocell = Eocathode - Eoanode where Eocathode and Eoanode are both standard reduction potentials.
Use the Nernst equation to predict the actual cell potentials. Show your work in the spaces provided including the
calculation of Q. Report your experimental potentials for comparison.
Nernst equation:
E cell  E ocell 
0.0592 V
log Q
n
(at 25oC; n is the number of electrons transferred in the balanced overall cell reaction)
Zinc-Copper cell:
Reduction half-reaction (cathode):
Eocathode = __________ V
Oxidation half-reaction (anode):
Eoanode = ___________ V
Overall cell reaction:
Eocell = ___________ V
Predicted Ecell
=
Experimental Ecell _______________
Zinc-Silver cell:
Reduction half-reaction (cathode):
Eocathode = __________ V
Oxidation half-reaction (anode):
Eoanode = ___________ V
Overall cell reaction:
Eocell = ___________ V
Predicted Ecell =
Experimental Ecell _______________
Copper-Silver cell:
Reduction half-reaction (cathode):
Eocathode = __________ V
Oxidation half-reaction (anode):
Eoanode = ___________ V
Overall cell reaction:
Eocell = ___________ V
Predicted Ecell =
Experimental Ecell _______________
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Part B: Batteries
Steps 1-8 (Lead storage battery)
You will observe 3 related processes:
- Initial charging of the battery
- Discharge of the battery. (Can it light a light bulb?)
- Recharging of the battery.
Charging of a battery is an example of electrolysis. An external power source (another battery or a power supply)
is used to force a non-spontaneous redox reaction to occur (you are stretching the spring!). Once the battery is
charged, the spontaneous reaction can be used to do work.
Steps 1-2 (Build the battery). Build the battery as described (see figure 3 on page 4 of the lab procedure).
Include the multimeter in the circuit. Show me your battery before you begin charging it.
Step 3 (Initial charging). Set the double throw switch to the charging position and turn on the power supply (set to
12 V). Watch what happens in the solution near each electrode. After charging for 5 minutes, set the double throw
switch to the neutral position and look for any color changes on the electrodes. Record your observations below.
On page 3 of the procedure, you will find two possible anode reactions that may occur during initial charging
(equations 4 and 5). The given potentials are standard oxidation potentials (Eoox). Copy the oxidation half-reactions
below and convert the given oxidation potentials to standard reduction potentials (Eored). (Simply change the sign
of the given potentials.)
Oxidation 1:
Eored =
Oxidation 2:
Eored =
Also on page 3, you will find two possible cathode reactions (equations 6 and 7). Copy those reductions and the Eo
for each (The given numbers are standard reduction potentials, Eored):
Reduction 1:
Eored =
Reduction 2:
Eored =
Charging of a battery is an example of electrolysis (using electrical energy to force a non-spontaneous redox
reaction to occur). Electrolysis is required when an electrochemical cell has a negative cell potential. What is the
minimum voltage setting on the power supply required to force all 4 reactions to occur? (Find the combination of
oxidation and reduction reactions that give the most negative cell potential.)
What were the bubbles that formed at the anode? What were the bubbles that formed at the cathode? On one of the
electrodes, a brown solid accumulated. Was this the anode or cathode? What is the brown substance?
Identity of bubbles at anode (oxidation) ___________
Identity of bubbles at cathode (reduction) ____________
Brown solid formed at the ____________ (enter anode or cathode)
Identity of brown substance _____________
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Steps 4 (potential). Set the multimeter on the 20 V setting. Set the double throw switch to the discharging
position. Record the voltage.
Measured potential after charging _________________ V
Step 5 (light bulb). Replace the multimeter with the light bulb or the motor is optional.
Did the light bulb or motor work? ____________
The overall reaction for the discharge of the battery is:
PbO2 (s) + Pb (s) + 2HSO4 (aq) + 2H+ (aq) → 2PbSO4 (s) + 2H2O (l)
You will find the two half reactions in the table on the next page. Write the two half reactions and show that you
get the correct overall (net) reaction. Calculate the standard cell potential from the standard reduction potentials.
Reduction (cathode):
Eored =
Oxidation (anode):
Eored =
_________________________________________________________________________
Net:
Eocell =
Use the Nernst equation to calculate the actual potential of this battery. Show your work including how you
calculate Q. [Write the formula for Q (what do you leave out?) and then plug in the appropriate concentrations.
Recall that the concentration of sulfuric acid is 3.6 M. What are the concentrations of HSO - and H+? HSO - is a
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weak acid so you can ignore the ionization of the second proton.]. “n” is the number of electrons that you crossed
out when you added the two half reactions.
Ecell = Eocell – (0.0592 V /n) log Q
Compare the calculated potential to the potential that you measured.
**Step 6 (recharging). Recharge the battery for 2 to 3 minutes and see if you can light the bulb or run the motor.
(**This is all you need to do for step 6. Ignore the rest of the instructions in the procedure.)
Recharging reverses the reactions that occur during discharging. (These are NOT the same as the reactions that
occur during the initial charging!) This requires energy which is provided by the power supply.
What is the minimum voltage setting required to recharge the battery? ______________ V
In addition to the reactions required to recharge the battery, another reaction occurs. This unwanted reaction wastes
a lot of energy. The unwanted reaction is the electrolysis of __________. (Hint: This problem was solved by the
invention of the “dry” cell.)
What is the minimal voltage required to do the initial charging of the battery? Why would this save energy? Why
might you want to use a higher voltage despite the fact that more energy would be wasted?
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