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					                                                                                                 95
Electrochemical Cells                                                         Experiment 15

1.0 Introduction
        You will need a table of electrode potentials today.
  In this experiment we will study the electrode reactions in electrochemical cells. We
will study primary voltaic cells, then electrolytic cells, and finally secondary
(rechargeable) cells. In voltaic (also called galvanic) cells, a spontaneous chemical
reaction occurs to generate electrical energy. In electrolytic cells, an external energy
source forces a nonspontaneous chemical reaction to occur. In all of these cell types,
oxidation occurs at the anode, while reduction occurs at the cathode. Electrons leave
the cell at the anode; electrons enter the cell at the cathode.
  In today’s experiment, the Procedure and Report Sheets are combined, so you can write
down your observations as soon as you make them. Please wash and return all metal
strips to the designated bottles.


2.0 Procedure
 In order to measure the cell voltages for the galvanic cells we will need to setup the
MeasureNet stations to measure voltage.
MeasureNet Task: Measure Voltage

              Your Actions:                                    MeasureNet Response:

           Press: MAIN MENU                             LCD: Shows Function Choices

           Press: F1 – Voltage                       LCD: Shows More Function Choices

         Press: F1 – Volt v Time                     LCD: Shows SELECT OPTION screen

             Press: DISPLAY                       LCD: Shows Voltage measurement screen

  Now the station will read voltage to the nearest millivolt on the screen when you connect
the voltage probe clips to a source.


2.1 Voltaic Cell, Primary Type                         Sn – Cu2+ cell
  Place about 25 mL of 0.5 M sodium sulfate, Na2SO4, solution in a 50 mL
beaker. Place a copper wire in the beaker with one end on the bottom and
the other end bent over the lip. (See the diagram, Figure 1, in the margin.)
  Place a piece of tin strip in the beaker and the other end bent over the lip
of the test tube. Drop, two or three crystals, of copper sulfate pentahydrate,
CuSO4 · 5 H2O into the beaker. Be careful to prevent the copper and tin
from touching and avoid unnecessary disturbance of materials in the
beaker.
 Carefully attach the cell to the MeasureNet station. The voltage leads are Figure 1
correctly connected when the red lead is connected to the  electrode of the
cell and the black lead is connected to the  electrode of the cell. This will result in a
positive voltage reading. With a positive voltage reading, electrons flow into the station

                                                                 College of San Mateo Chemistry Dept
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Electrochemical Cells                                                       Experiment 15


 from the black terminal, through the station, and then out through the red terminal.
Record the orientation of the leads and the maximum voltage for the cell on the report
sheet.

 1.   Maximum cell voltage:                       ______________________________

 2.   Write the equation for the half-cell
      reaction taking place on the Sn strip:      ___________________________________

 3.   Write the equation for the half-cell
      reaction taking place on the Cu wire:       ___________________________________

 4.   The copper is connected to the (   ) lead of the Station.

 5.   Electrons ( enter leave ) the copper wire at its terminal.

 6.   The Sn strip is the ( positive negative ) terminal.

 7.   ( Positive negative ) ions drift towards the Sn strip.

 8.   The Sn strip is the ( anode cathode ).

 9.   ( Oxidation reduction ) occurs at the copper wire, which is the ( cathode anode ).

10. Combine the half-cell equations in 1 and 2 to obtain the overall equation.

      ________________________________________________________________________

  Using a table of reduction potentials calculate the expected voltage for this cell. What
is the percent error in your reading? Is
there any reason why the reading might
                                               11. Calculated Cell voltage
not be expected to be exactly the same as
the calculated voltage?
                                               12. Percent Error


                     2.2   Electrolytic Cells
                       Electrical energy can cause oxidation and reduction of either solute
                     particles, or water solvent, or the cell electrodes themselves. Inert
                     electrodes such as carbon are often used to prevent the electrode
                     materials from interfering with the redox reactions under study. Be
                     sure that the carbon electrodes never touch when the power is turned
                     on.
                     Electrolysis of a Na2SO4 Solution
                       To make the apparatus shown at the left, place 35 mL of 0.5 M
                     Na2SO4 solution in a 50 mL beaker. Support the carbon electrodes in
                     the solution using a ceramic ring clamp and separate the electrodes as
      Figure 2       far as possible.

                                                               College of San Mateo Chemistry Dept
                                                                                               97
Electrochemical Cells                                                       Experiment 15


   Decide which electrode will be the anode and which the cathode. With this in mind,
connect the electrodes to the power supply. Turn on the power supply and let the current
flow for 2-3 minutes.


13. Observation



14. Associate the relative amounts of products at the electrodes.




15. Test for the presence of H3O+ (aq) and OH- (aq) produced during reaction by
    touching both red and blue litmus papers, in turn, against the submerged
    surfaces of the carbon electrodes.

           Anode:          ________________         Cathode:         __________________



     In this cell, the anode is the terminal at which electrons were withdrawn from
     the cell by the positive terminal of the power supply. With this in mind, answer
     the following questions.


16. Particles at the anode available for reaction are _______________ ions,
    _____________ ions, and ______________ molecules.


17. During the electrolysis, a gas was liberated at the anode and the litmus turned (
    red blue ) indicating an increase in ______________ ions. The gas was
    ______________.

18. Therefore, the ( oxidation reduction ) of H2O taking place at the anode results
    in the formation of ______________ and ______________.


19. Assuming the SO42- is too stable to react, write the equation for the half-reaction
    at the anode:



     _______________________________________________________________________________




                                                               College of San Mateo Chemistry Dept
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Electrochemical Cells                                                       Experiment 15



     In this cell, the cathode is the terminal receiving electrons from the negative
     terminal of the power supply. With this in mind, answer the following
     questions.


20. Particles at the cathode available for reaction are _______________ ions,
    _____________ ions, and ______________ molecules.


21. Evidence for a chemical change at the cathode was liberation of a gas, and the
    litmus turned ( red blue ) indicating an increase in ______________ ions. The
    gas was ______________.

22. Therefore, the ( oxidation reduction ) of H2O taking place at the cathode
    results in the formation of ______________ and ______________.


23. Write the equation for the half-reaction occurring at the cathode:

      ______________________________________________________________________________


24. Combine the two half-reaction equations to arrive at the over-all cell reaction
    equation for the electrolysis of Na2SO4 solution.

     _______________________________________________________________________________



25. Do the relative amounts of gases generated during the electrolysis correspond to
    the coefficients of the gas products in your over-all cell reaction equation?
    Explain.




                                                               College of San Mateo Chemistry Dept
                                                                                              99
Electrochemical Cells                                                      Experiment 15



2.3 Electrolysis of a CuSO4 Solution (Electroplating)
  Electroplating is the process in which an object is made the cathode of an electrolytic
cell, and a protective layer is plated, or coated, upon the object. In this section, the
electrodes are not inert. An iron nail is used for the cathode, and a copper wire is used
for the anode. Copper sulfate is used for the electrolytic solution.
  Before carrying out this part of the experiment, try to make some reasonable
predictions about the reactions that might take place.


26. The particles reduced at the nail cathode are most likely to be ( Fe Cu2+ ).

27. The particles oxidized at the wire anode are most likely to be ( Cu Cu 2+ ).

28. Using an electrode potentials table, explain why H2O (in a neutral solution) would
    not be reduced.



29. Using an electrode potentials table, explain why H2O (in a neutral solution) would
    not be oxidized.



30. Write a “possible” half-reaction
    equation involving copper at the    ______________________________________________
    anode.

31. Write a “possible” half-reaction
    equation involving copper at the    ______________________________________________
    cathode.

32. If the nail is to be the cathode, the (   ) terminal of the power supply must be
    attached to it.




                                                              College of San Mateo Chemistry Dept
100
Electrochemical Cells                                                     Experiment 15


  Now you will build an electroplating cell (Figure 3, right). Place 35–40
mL of a CuSO4 solution into a 50 mL beaker. Attach the leads from a
power supply to an iron nail and a short length of copper wire. Use your
answers to Questions 26 & 27 to help you attach the leads. Making sure
they don’t touch, set the nail and wire into the CuSO 4 solution. Also,
don’t allow the clips on the power supply leads to touch the solution.
Turn on the power, and watch for a few minutes. Unplug the power
supply, and observe the nail.

33. What evidence of chemical change do you observe?


34. What is deposited on the nail?
                                                                                   Figure 3
35. Write the anode half-
    reaction equation.        _____________________________________

36. Write the cathode half-
    reaction equation.          ____________________________________

37. What would eventually happen to the wire if the power supply
    were left on continuously?


38. Label in Figure 3 the anode and the cathode, and the sign of
    each electrode. Also show the path of electrons through the
    cell and external circuit.

   Reverse the leads on the power supply so the iron nail is the
anode in the circuit and the copper wire is the cathode. Turn on
the power, watching for a few minutes.

39. What evidence of chemical change do you observe?


40. Write the anode half-
    reaction equation.        __________________________________
41. Write the cathode half-
    reaction equation.          _________________________________
                                                                                    Figure 4
42. What would eventually happen to the wire if the power
    supply were left on continuously?

43. Label in Fig. 4 the anode and cathode, and the sign of each
    after the terminals were reversed. Also show the path of the
    electrons through the cell and external circuit.



                                                             College of San Mateo Chemistry Dept
                                                                                             101
Electrochemical Cells                                                      Experiment 15




2.4 Lead Storage Cell (Secondary Type)
  The voltaic cell studied in Section 2.1 is called a primary cell. A primary
cell produces an electric current, but is not rechargeable. A secondary cell,
or storage cell, is a cell whose energy can be renewed by passing an electric
current through it in the direction opposite to the current flow generated by
the cell itself.
  A secondary cell behaves as a voltaic cell as it discharges (to produce an
electric current) and as an electrolytic cell as it charges (accepting an
electric current from an external power source). To see in chemical terms
how this can occur, ponder the following cell reaction:
           reactants       products + electrical energy
                                                                                 Figure 5
  If all reactants and products remain in electrical contact, the equilibrium is
preserved. The discharge reaction goes from left to right as written, to generate an
electric current; the charge reaction, occurring only in the presence of an external
electric current, goes from right to left as written.
  Above on the right is a diagram (Figure 5) of a discharged lead storage cell, much like
a single compartment in a car battery. The electrodes are first connected to a power
supply and the power supply turned on to charge the cell. The electron flow is given
by the arrow on the diagram. The reaction at each electrode involves the lead sulfate,
which is formed on the surface of the lead (that’s why car batteries are so heavy!)
electrode as soon as the lead touches the sulfuric acid. Answer Questions 44 and 45
on the report sheet before proceeding.

   44. If PbSO4 gains electrons, the possible products are SO42- and
       _____________________________.

   45. If PbSO4 loses electrons, the possible products are SO42- and
       _____________________________.




Charging the Lead Storage Cell (Clean the Pb with steel wool first.)

  Half-fill the small lead cells with 3 M H2SO4. Connect the electrodes
to a power supply; allow the current to flow for 3–5 minutes. Do not
let the electrodes touch each other. Notice that since our power
supplies produce a 5 V current there will also be gas formed in this
cell. When auto batteries are charged this is prevented by carefully
controlling the power supply voltage. For the following answers ignore
the gas formation reactions.

                                                            Figure 6


                                                              College of San Mateo Chemistry Dept
102
Electrochemical Cells                                                     Experiment 15




46. A ______________ colored solid has deposited on the electrode connected to the positive
    terminal of the power supply. This is the result of ( oxidation reduction ).
47. The anode half-reaction equation is:
         PbSO4 + 2 H2O  PbO2 + ______________ + ______________ + ______________
48. The cathode half-reaction equation is:
                        PbSO4 + 2 e-  ______________ + ______________
49. The single equation for the complete reaction during charging of the cell is:
    ______________________________________________________________________________________
50. Connect the cell to the voltage probe and record the maximum voltage_________________ v
   As you read the voltage, the lead cell was discharging. As the cell was discharging :
                                                                                        
51. Electrons ( enter leave ) the pure Pb electrode.
52. The PbO2 electrode is the ( anode cathode ).
53. The external terminal of the PbO2 electrode is ( positive negative ).
54. The anode half-reaction           __________________________________________________
    equation is:
55. The cathode half-reaction         __________________________________________________
    equation is:
56. The single equation for the complete reaction during discharging of the cell is:
    ______________________________________________________________________________________
57. How do the equations in 49 and 56         __________________________________________
    compare?
58. In which direction(s) do the SO42- move during       _______________________________
    discharge?
59. Complete Figure 6 on previous page, showing a charged cell attached to a station. Show
    direction of e- flow during discharge, anode and cathode,  and  external terminals and
    motion of ions during discharge.
60. Using an electrode potential table, calculate the E for this cell ___________________ v
61. Calculate per cent error between 50 and 60. _____________________ %




                                                             College of San Mateo Chemistry Dept

				
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