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ELECTROCHEMICAL CELLS

 

Prelab

 

NAME:________________________________________________ PERIOD:_______

 

1. Given the reaction:  Al(s)  +  Cu+2(aq) ®  Al+3(aq)   +   Cu(s)

 

a. What is being oxidized?_________

b. What is being reduced?_________

c. What is the oxidizing agent?_________

d. What is the reducing agent?_________

 

e. Write the half-cell reaction for the oxidation.

 

 

f. What is the standard voltage for this half-cell reaction?_________

 

 

g. Write the half-cell reaction for the reduction.

 

h. What is the standard voltage for this half-cell reaction?_________

 

 

i. Write the balanced equation for the cell reaction.

 

 

j.  What is the standard cell voltage for the reaction?_________

 

 

k. Is this reaction spontaneous under standard conditions?_________

 

2. Given the reaction:  Ag (s)   +   H+1(aq) ®  Ag+1(aq)   +   H2 (g)

 

a. Write the balanced equation for the cell reaction.

 

b.  What is the standard cell voltage for the reaction?_________

 

 

c. Is this reaction spontaneous under standard conditions?_________

 

d. Is this in agreement with what you would predict based on the position of the Ag on the Activity Series for metals and its ability to replace hydrogen? Explain your answer.

 


ELECTROCHEMICAL CELLS

 

            The potential difference or voltage of an electrochemical cell depends on the difference in the electron attracting ability of the two half-cells. The half-cell with the greater electron attracting ability will be the reduction half-cell and is designated the cathode. The half-cell with the lower electron attracting ability will be the oxidation half-cell and is designated the anode. Electrons leave the electrochemical cell from the anode, flow through the external circuit, and enter the electrochemical cell at the cathode. The external circuit may contain a voltmeter to measure voltage and an ammeter to measure current. The circuit is completed using a salt bridge, which maintains electrical neutrality in the two half-cells. Ions in the salt bridge migrate into the two half-cells to balance the charge of any ions lost or produced in the half-cells.

            A diagram of a typical electrochemical cell is shown below.

 

External Circuit

 

Zn Electrode

 

Cu Electrode

 

 

1.00M Cu(NO3)2

 

1.00M

Zn(NO3)2

 
 

 

 

 

 

 

 

 

 

 

 

 

 

 

Cathode Reaction

Cu+2(aq)  +  2 e-   ®   Cu(s)

 

Anode Reaction

Zn(s)   ®   Zn+2(aq)  +  2 e-

 
 

 

 


In a standard cell, the concentration of the dissolved compounds is 1.00M. A table the standard reduction potentials provides the half-cell potentials:

            Cu+2(aq)  +  2 e-   ®   Cu(s)                    Eo = +0.34V

            Zn+2(aq)  +  2 e-   ®   Zn(s)                     Eo = -0.76V

Under standard conditions, the more positive half-cell has the greater electron attracting ability and will be the reduction and the cathode reaction. The less positive half-cell has the lower electron attracting ability and will be the oxidation and the anode reaction. The anode reaction is reversed and the sign of the half-cell voltage changed. Each half-cell reaction is multiplied by an integer so the number of electrons lost equals the number of electrons gained. These half-cell reactions are added to give the overall balanced equation. The half-cell voltages are not multiplied by the integers but are simply added algebraically.

            Cathode           Cu+2(aq)  +  2 e-   ®   Cu(s)                                Eo = +0.34V

            Anode              Zn(s)   ®   Zn+2(aq)  +  2 e-                                 Eo = +0.76V

            Overall Cu+2(aq)  + Zn(s) ®   Cu(s)  + Zn+2(aq)                   Eocell = +1.10V

In the cell diagrammed above:

a.       In which direction do the electrons flow in the external circuit? Why?

b.      What happens to the mass of the copper electrode? Why?

c.       What happens to the mass of the zinc electrode? Why?

d.      What happens to the concentration of the Cu+2 ions in solution? Why?

e.       What happens to the concentration of the Zn+2 ions in solution? Why?

            The salt bridge is a critical component of the electrochemical cell since it provides ions that migrate into the two half-cell compartments to maintain electrical (charge) neutrality. As the cathode reaction occurs, there will be a surplus of NO3-1 ions in that compartment. K+1 ions migrate out of the salt bridge to balance these NO3-1 ions. As the anode reaction occurs, there will be a surplus of Zn+2 ions in that compartment. NO3-1 ions migrate out of the salt bridge to balance these Zn+2 ions. A very water-soluble ionic compound that will not react with the compounds in either half-cell is chosen for the salt bridge. NaNO3 and NH4NO3 are also used. A high concentration of ions is used in the salt bridge so that the salt bridge is not easily depleted of ions otherwise the ions from the two half-cells will start to migrate through the bridge.

            On a voltmeter, the black terminal is the negative terminal and the terminal where electrons leave the electrochemical cell and enter the meter. If the voltage on the meter reads a positive value when the black terminal is connected to an electrode, that electrode is the anode since the anode half-cell involves oxidation and releases electrons to the external circuit. The red terminal is the positive terminal and the terminal where electrons leave the meter and enter the electrochemical cell. If the voltage on the meter reads a positive value when the red terminal is connected to an electrode, that electrode is the cathode since the cathode half-cell involves reduction and gains electrons from the external circuit. You will sometimes hear the expression, reduce-red-cats, to help you remember the connection.

 

In this experiment, you will:

a.       Create a series of electrochemical cells.

b.      Measure the voltage of the cells.

c.       Determine which half-cell is the anode and cathode.

d.      Determine the order in the reduction potentials of the half-cells.

e.       Compare the experimental voltages to the theoretical voltages calculated from a table of standard reduction potentials.

 

Procedure:

 

1. Obtain a 9cm piece of filter paper that has been cut to make a cell template.

 

2. Write on each sector the atomic symbols of the metals: Ag, Zn, Cu, and Pb.

 

3. Place the filter paper on a plastic sheet and place a piece of the metal on the corresponding

    section of the paper.

 

4. Place 2 drops of a 1M solution of the corresponding metal ion (AgNO3, Zn(NO3)2, Cu(NO3)2,

    Pb(NO3)2 solution on the paper at the edge of each metal piece so that each metal is in contact

    with its solution.

 

5. Drop 1M KNO3 into the middle of the filter paper to form a salt bridge. Be sure that the

    KNO3 solution soaks outward and contacts all the other solutions.

 

6. Use a voltmeter and measure the voltage difference between every combination of two

    different metals. Touch the probe to each combination of metals. Be sure make good contact

    with the metal. Arrange the probes to achieve positive voltage values.

 

7. Record your cell voltages.

 

8. Decide which metal is the cathode and which metal is the anode.

 

9. Calculate the theoretical voltages for each cell.

 

10. Arrange the metals in order of strength as a reducing agent from high to low. 

 

CBL Instructions Using the CBL in the Multimeter Mode:

 

1. Plug the voltage probe into Channel 1 of the CBL.

 

2. Attach a short piece of Nichrome wire to each lead.

 

3. Turn on the CBL unit and press the MODE button to switch the CBL to Multimeter mode. It

     should be reading voltage. Check to see if there is a mV or V to the right of the center display.

     Check to see if Channel 1 and V are displayed in the upper left corner.

 

4. Touch the Nichrome wire to each combination of metals. Be sure make good contact with the

     metal.

 

5. When the meter gives a steady reading, record the voltage from the CBL.

 

6. Decide which metal is the cathode and which metal is the anode.

 

7. Calculate the theoretical voltages for each cell.

 

8. Arrange the metals in order of strength as a reducing agent from high to low.

 

II. CBL Instructions Using the CHEMBIO Program:

 

1.   Prepare the voltage probe.

a. Plug the voltage probe into Channel 1 of the CBL.

b. Attach a short piece of Nichrome wire to each lead.

c. Connect the CBL System to the TI-82 calculator using the unit to unit link cable. Push the

    cable in firmly at both ends.

 

2.   Turn on the CBL unit and calculator. Press [PRGM] and select CHEMBIO. Press [ENTER]. (CHEMBIO may be under the APPS menu on the TI-83+). “Prgm CHEMBIO” appears. Press [ENTER].  “ Vernier Software-Biology and Chemistry with the CBL” appears. Press [ENTER] again to go to the MAIN MENU. If the CBL and the calculator are not turned on and the link cable pushed in firmly at both ends, the message “ Link Error” will appear. Be sure that the link cable is firmly pushed into the CBL and calculator and press the On/Halt button on the CBL. Press [ENTER] again to go to the MAIN MENU.

 

3.   Set up the calculator and CBL for the voltage probe.

a. Select SET UP PROBES from the MAIN MENU. Press [ENTER].

b. Enter “1” as the number of probes. Press [ENTER].

c. Select VOLTAGE from the SELECT PROBE menu. Press [ENTER].

d. Enter “1” as the channel number. Press [ENTER].

 

4.   Set up the calculator and CBL for data collection.

a. Select COLLECT DATA from the MAIN MENU. Press [ENTER].

b. Select MONITOR INPUT from the DATA COLLECTION menu. Press [ENTER].

     No data will be stored.

c. Prepare the filter paper and metals as directed above.

      d. Touch the Nichrome wire to each combination of metals. Be sure make good contact with

          the metal.

e. When the meter gives a steady reading, record the voltage from the CBL.

f. If the voltage is negative, reverse the leads and record the voltage.

g. Decide which electrode is the cathode and anode.

h. Calculate the theoretical voltages for each cell.

i. Arrange the metals in order of strength as a reducing agent from high to low.

j. Press + to QUIT monitoring data. Select QUIT to exit the program.

 

 

This experiment is adapted from A Small-Scale Electrochemical Cell and An Electrochemical Series from Cell Data in Chemtrek by Stephen Thompson, Allyn and Bacon, 1990, p. 311-315.


 

ELECTROCHEMICAL CELLS

 

NAME:__________________________________________ COURSE:___________

 

LAB PARTNER:___________________________________ PERIOD:___________

 

DATA TABLE

 

Cell

Cathode

Anode

Experimental Voltage

Theoretical Voltage

Zn-Pb

 

 

 

 

Zn-Cu

 

 

 

 

Zn-Ag

 

 

 

 

Pb-Cu

 

 

 

 

Pb-Ag

 

 

 

 

Ag-Cu

 

 

 

 

 

Calculations:

For each of the cells above: write the oxidation half-cell reaction with its standard voltage, write the reduction half-cell reaction with its standard voltage, write the balanced cell reaction with the cell voltage. Include the physical states.

 

 

 

 

 

 

 

 

 

 

 

 

 

 


Questions:

1. Why do your experimental voltages differ from the theoretical voltages?

 

 

 

 

 

 

 

2. Arrange the metals in order of strength as a reducing agent from high to low (left to right).

 

 

 

 

 

 

 

 

 

3. What is the relationship between the strength of the metal as a reducing agent and its standard

     reduction potential (be specific about the sign of the standard reduction potential)?

 

 

 

 

 


 

ELECTROCHEMICAL CELLS

 

SETUP SHEET

 

10        Vials with a small piece of Zn, Pb, Ag, Cu

10        Sheets of plastic (overhead transparency sheets will work)

10        Sets of Berol dropper bulbs filled with 1M solutions of:

             Zn(NO3)2, Pb(NO3)2, AgNO3, Cu(NO3)2, KNO3

10        Wooden racks to hold the droppers

10        Boxes Kimwipes

70        9 cm filter paper cut in the form of the template.

10             CBL with voltage probes and link cables

10             CBL voltage adapters

20             Short pieces of Nichrome wire

 

 

 

 

 

 

 

 

 

 

 

 

 


100ml   1M Zn(NO3)2, Pb(NO3)2, AgNO3, Cu(NO3)2, KNO3

 

 


ELECTROCHEMICAL CELLS

 

Instructor’s Notes

 

Cell

Cathode

Anode

Experimental Voltage

Theoretical Voltage

Zn-Pb

Pb

Zn

+0.442

+0.63

Zn-Cu

Cu

Zn

+0.984

+1.10

Zn-Ag

Ag

Zn

+1.40

+1.56

Pb-Cu

Cu

Pb

+0.624

+0.47

Pb-Ag

Ag

Pb

+0.907

+0.93

Ag-Cu

Ag

Cu

+0.401

+0.46

 

 

1. Good contact with the metals is important. The student may need to add an additional drop of the ion solution if the paper starts to dry out. Buffing off the oxide coating on the metals will help. Use flats pieces of zinc and copper plate rather than zinc and copper shot. Flatten the silver wire so that it will make better contact with the solution.

 

2. Do not add so much KNO3 that it starts to dilute the other solutions. Do not add so much of the metal ion solutions that they run into the other sections.

 

3. Analog or digital volt meters or multimeters can be used. Use the smallest scale that will cover 0-2V.

 

4. If a full size cell is set up as a demonstration, a salt bridge can be made using a rolled-up piece of filter paper soaked in KNO3 and folded so it makes contact with the solutions.