Measuring Electrochemical Cells Download

Transcript
Electrochemical Cells
Determining Relative Reduction Cell Potentials
In lab there are five half-reactions. By measuring the voltage between a variety of halfcells you are to place the half-cell reductions in order of decreasing half-cell potential.
The following cells can be made and measured in the 24-well plate.
Ag+ + e → Ag(s)
Zn2+ + 2e → Zn(s)
Fe3+ + e → Fe2+(s)
Cu2+ + 2e → Cu(s)
2I2 +2e → I(aq)
Create the five half-cells above in the well plate. Fill five wells half-way with a different
ion solution. In these five solutions place the appropriate metal or graphite electrode.
For the Fe3+ and I solutions use a graphite electrode. The solutions in the well plate
should not go down the drain. When it is time to clean up your cells, empty the solutions
into the metal waste beaker in the hood.
Sign Convention of the Volt Meter
The volt meter (truly a multi-meter since it also measures current and impedance) that
you have should be turned on by setting the center dial to the 2V DC position. (All
electrochemical cells generate a continuous flow of electrons, while the voltage from the
wall provides an oscillating current flow, AC). The volt meter will read the potential
difference between two half-cells. The potential will read positive when the red lead (+)
is attached to the cathode (reduction half-cell) and will read a negative voltage when the
red lead is attached to the cell anode (oxidation half-cell). Naturally, switching the leads
will change the sign of the voltage on the meter. The sign of the voltage is useful for
determining which cell is at the higher reduction potential.
Filter Paper Salt Bridges
You will need multiple salt bridges. It is important that each combination of cells you
test have its own salt bridge. With a pencil, label each salt bridge indicating which end is
placed in which solution. After labeling each salt bridge, place the collection on a watch
glass and drip saturated KNO3(aq) solution over the strips so that each is completely
dampened.
Measuring Cell Voltages
Select a pair of half-cells to measure the potential difference between. Pick one on your
own or begin with zinc and iodine cell combination. Connect one of the leads from the
volt meter to the zinc electrode and place it in the zinc ion solution. Attach the other volt
meter lead to the graphite lead and place the graphite in the iodine solution. Find the salt
bridge that you have labeled Zn and I2 . Place the salt bridge into each half cell. Record
the voltage of the half cell, and identify the sign of the voltage. Be sure that from the
sign of the voltage you can identify which cell is being oxidized and which cell is being
reduced. Does Zn oxidize I2 or does Zn reduce I2?
For each cell you measure, write the cell notation in your lab notebook. For the case of
the zinc /iodine cell you should write:
Zn(s)│Zn2+(1M)││I2(aq 1M),I(1M) │C(graphite)
The electrodes are identified in the cell notation at the beginning (anode) and at the end
(cathode). Notice that the reduction of I2 to I all takes place in the aqueous phase so
there is not a phase separation bar, │, between reactants and products.
From this measurement you should conclude that 2I + 2e → I2(aq) is a greater half-cell
voltage than Zn2+ + 2e → Zn(s). If you determine a half-cell voltage lower than zinc,
you do not need to measure this half-cell against I2 because if it is lower than Zn it is also
lower than the I2 reduction potential.
Continue measuring cell potential differences for different half cell combinations, writing
the cell notation and measured voltage until you can order the five half-cell reactions
above from largest to smallest reduction potential.
Concentration Cell
Choose a copper unknown solution to investigate. Construct the following
electrochemical cell and record the voltage observed:
Cu(s)│Cu2+(???M)││Cu (1 M)│Cu(s)
From the sign of the observed voltage in your concentration cell, is the 1M Cu2+ solution
the anode or the cathode?
Use the Nernst equation to calculate the Unknown concentration:
0
ECell  ECell

RT
ln Q
nF
where R= 8.314 J/mol K; n = # of electrons transferred
F = 96, 500 J/Vmol
Electrolytic Cell
At the front of the room your TA has an electrolytic cell. This is a non-spontaneous
redox reaction which when attached to a battery of greater voltage can be made to run
spontaneously.
The non-spontaneous reaction in the electrolytic cell is probably pretty complex because
we are using a very large potential, but one redox reaction occurring is:
H2O  2e– + 2H+ + ½O2
2H2O +2e–  H2 + 2OH–
Eºcell = 1.23V
Eºcell = 0.83V
Using the above reaction answer in your notebook as you run this cell with your teaching
assistant:
1. At which electrode (anode or cathode) will hydrogen gas be produced?
2. If phenolphthalein is added to the cell, which electrode (anode or cathode) will turn
pink?
Cells in Series
What is the voltage of a car battery? 12 volts. How does the battery get that large of a
voltage? The battery is really six 2 volt batteries wired in series. In this experiment
prepare a battery set in series.
At the front of the room there are two Daniels cells Zn(s) + Cu2+  Zn2+ + Cu(s).
Measure and record the cell voltage for each cell.
Now connect the two cells in series as is shown in the figure below.
salt bridge
Zn
Cu
Zn
Cu
2.2 V
+
Measure and record the voltage of the cells in series. Examine the 9-volt battery at the
front of the room. Estimate the voltage of a single cell in the 9-volt battery.