LAB 28 DEMO: ELECTROCHEMICAL CELLS
BACKGROUND:
You learned in lab 27 metal ions are reduced by other metals. That was a qualitative lab, you only determined which ion was a better thief of electrons, not how much better. This lab will measure how good a thief a metal ion is. You also had a spontaneous reaction occur because you placed the metal directly into the metallic ion solution. In this lab you will keep the two half-reactions separate from each other and will measure the flow of electrons through a wire and ultimately through a voltmeter.
Now you know which ion is the best thief in each set up, you can just refer back to your results in lab 27. Knowing which metallic ion steals electrons in each set up will allow you to “see” which way electrons flow through the wire. The needle on the voltmeter will also point in that direction. You also know reduction is occurring in that cell. The metal strip in this cell will be called the cathode. The metal strip in the other cell will have oxidation occurring at it and will be called the anode. For any electrochemical cell to work, a complete circuit is needed. To complete this circuit a salt bridge will be used. Your instructor will discuss how it works and what its made of. You must remember that electrons do not swim, only ions do.
PROCEDURE:
Build the following set-up:
v
Zn salt bridgeCu
Zn(NO3)2 Cu(NO3)2
PART 1
Zn(s) / Zn+2(aq) and Cu(s) / Cu+2(aq)
Record the voltmeter reading from the demonstration when:
A. no salt bridge in place
B. salt bridge in place
PART 2
Pb(s) / Pb+2(aq) and Cu(s) / Cu+2(aq)
Record the voltmeter reading from the demonstration when the salt bridge is in place.
PART 3
Ag(s) / Ag+(aq) and Cu(s) / Cu+2(aq)
Record the voltmeter reading from the demonstration when the salt bridge is in place.
If the needle is deflected in the wrong direction what does that mean?
PART 4
effect of changing concentration on voltage
1. Cell will be set up as in part 1 except a porous cup will be used instead of the salt bridge.
2. Inside the porous cup will be the zinc strip and Zn(NO3)2 solution. The copper strip and Cu(NO3)2 solution will be in the beaker.
The porous cup allows ions to move through it but the water molecules cannot.
3. Connect the leads as before and note the reading. If it is higher than part 1 it is because there is less resistance through the cup than the bridge and therefore ions can move more freely and more electrons flow, thus increasing the voltage. Note the voltage.
4. Now the instructor adds 2 M Na2S slowly to the beaker, note the voltage. They will continue to add the sodium sulfide until all change is noted.
Questions and Calculations:
1. Why won’t a cell operate without the salt bridge?
2. What is the purpose of the salt bridge or porous
cup?
3. A. Write the half-reactions occurring in each cell
in part 1.
B. Write the overall reaction.
C. What is the flow of electrons through the wire?
D. What is the direction of movement of ions
through the salt bridge?
E. Using the table of Eo values in your packet,
determine the expected value. Explain why
experimental values are different than table
values.
4. Write the overall reaction for step 4 and explain the
results obtained using Le Chatlier’s principle.
5. Sketch a cell set up similar to the ones used in steps
1-3 using the zinc half-cell in one and Br2(l) for the other. Since bromine is a liquid an inert electrode must be made of a conducting metal, use platinum.
A. show the direction of electron flow through
the wire.
B. write the equation for each half-reaction
occurring at each electrode and the overall
reaction.
C. predict the voltage of the cell. (Eo)
D. It is known that heat is produced when the
reaction occurs. Would the voltage of the cell increase or decrease if the temperature was increased? Explain using Le Chatlier’s principle.