Oxidation-Reduction Reactions and Voltaic Cells

 

Purpose – In this experiment you will compare the reactivity of four different metals and their ions.  Then you will build voltaic cells out of these metals and their ions and compare the voltage of the different cells.  Then you will see how the change in concentration of an ionic metal ion changes the voltage of the cell.

 

I.        Oxidation-Reduction Reactions of Metal Atoms and Ions

 

                    A. Clean and rinse with distilled water, four small test tubes.  Shake excess water out.  Fill each tube half full of Zn(NO3)2 solution.  Clean strips of Zn, Cu, and Ag and Sn metals by gently sanding the bottom one-inch strip.  Submerge the cleaned end of the strips into the zinc solution for several minutes.  Record your observations in the your notebook in a similar table to the one below.

 

                    B. Rinse the test tubes and clean the metal strips.  Fill the tubes half full of Cu(NO3)2 solution and place the metal strips in the solutions.  Record your observations.

 

                    C. Repeat the above with AgNO3 and Sn(NO3)2 solutions.

 

 

Zn (s)

Cu (s)

Ag (s)

Sn (s)

Zn2+

 

 

 

 

Cu2+

 

 

 

 

Ag+

 

 

 

 

Sn2+

 

 

 

 

 

Q1. Write balanced chemical equations for each system tested.  (If you observed no change in the system write NR, for NO REACTION to the right of it).  Pick two of the reactions and explain why you chose the products you did.

 

Q2. What patters are shown in these data?  (Hint: compare the relative reactivity of the metals to each other.  Compare the relative reactivity of the metals ions to each other.  Identify connections between metals and ions reactivity.)

 

Q3. Pick one of the reactions and draw a picture that shows how the metal atoms and ions interact.  Explain in words how your picture illustrates your observations.

 

II.       Voltaic Cell EMF

 

          A. Obtain a U-tube with a length of string inside it that is soaked with ammonium nitrate, potassium nitrate, or some-other strong electrolyte.  This is our salt bridge.

 

          B. Clean, rinse with distilled water, and shake dry four test tubes.  Label each tube “Zn,” “Sn,” “Ag,” and “Cu”.  Fill the tubes with Zn(NO3)2, Sn(NO3)2, AgNO3, and Cu(NO3)2.  Put a piece of Zn metal into the zinc nitrate solution, a piece of tin into the tin (II) nitrate solution, and so forth.  Bend the metal strip over the edges of the test tube and store in a 100 mL beaker.

 

          C. Assemble the Zn/Sn cell.  Rinse (by dipping the ends of the salt bridge into water) and blot dry. Place the two test tubes into a 100 mL beaker and place one leg of the salt bridge into each test tube.  Clip one wire to the Zn metal (in the zinc half cell) and connect the other end to a voltmeter set to measure 0-2 Volts DC.  A second wire will run from the voltmeter to the Sn metal.  After the voltmeter has stabilized, record the absolute value of the voltage.  Disconnect the wires, reconnect and re-measure the voltage.  You have just measured the voltage in the Zn | Zn2+¦ Sn2+  | Sn cell.

 

          D. Assemble and record the voltage from each of the possible six cells from the four half-cells.  Rinse and blot dry the ends of the salt bridge between each reading.  Save the Ag+ and Cu2+ solutions for the following experiment.

 

      

 

Go through the process again obtaining a second reading and average the two.  Record the data in your notebook. 

 

Cell = ½ cell (1) + ½ cell (2)

First Voltage

Second

Average

1. Zn | Zn2+¦ Sn2+  | Sn

 

 

 

2. Zn | Zn2+¦ Ag+  | Ag

 

 

 

3. Zn | Zn2+¦ Cu2+  | Cu

 

 

 

4. Cu | Cu2+¦ Sn+  | Sn

 

 

 

5. Cu | Cu2+¦ Ag+  | Ag

 

 

 

6. Sn | Sn2+¦ Ag+  | Ag

 

 

 

 

Q4.  What patterns exist in the voltage readings for the different combinations of half-cells?  How do these patterns compare to those found in Q2?

 

Q5.  What relationships are there among the voltages of the Sn | Sn2+|| Ag+  | Ag, Zn | Zn2+|| Sn2+  | Sn,  and Zn | Zn2+|| Ag+  | Ag?

 

III.       Concentration Effect in Voltaic Cell EMF

 

          A. The Ag | Ag+ and Cu | Cu2+ half-cells will be used in this experiment.  You will need the following Ag+ molarities: 0.2 0 M, 0.020 M, 0.0020 M and .00020 M.  The last three can be prepared by a technique called serial dilution.  Use your burets to prepare the following solutions.  The 0.020 M solution is prepared by diluting 1.00 mL of the 0.20 M Ag+ solution with 9.00 mL of distilled water.  The 0.0020 M solution is prepared by taking 1.00 mL of the 0.020 M Ag+ solution and adding 9.0 mL of DI water.  In a similar manor, the 0.0002 M can be prepared.  Place a few milliliters of each Ag+ solution in separate, clean, dried, labeled small test tubes.

 

          B. Place a cleaned Ag strip in the 0.00020 Ag+ solution and connect this half-cell to the Cu half-cell.  Measure the voltage as described above and record in your notebook.

 

Remove the Ag strip and salt bridge, rinse both and construct a cell using the 0.0020 M Ag+ solution.  Repeat this procedure with the 0.020 M and the .2 M solutions.

 

Obtain a second reading of the voltage for each cell.  Record these and compute an average reading for each cell.

 

Cell

First Voltage Reading

Second

Average

Ag | 0.0002 M

Ag+¦ Cu+2  | Cu

 

 

 

Ag | 0.002 M

Ag+¦ Cu+2  | Cu

 

 

 

Ag | 0.02 M

Ag+¦ Cu+2  | Cu

 

 

 

Ag | 0.20 M

Ag+¦ Cu+2  | Cu

 

 

 

 

 

   Q6.  Plot the above data and the theoretical results calculated using the Nerst equation.  You might want to plot the concentration data on a log scale.  You can plot the two data sets on the same graph by entering your data into Excel as follows.

 

Cu2+ concentration Average Exp. Voltage      Theoretical Voltage

0.00020                  V1 (exp.)                                     V1

0.0020                    V2  (exp.)                                    V2

etc.