Chemistry 115B

Galvanic (Voltaic) Cells and Electrode Potential

Oxidation-reduction (redox) reactions occur when electrons are given up by the substance being oxidized (the reducing agent) and simultaneously gained by the substance being reduced (the oxidizing agent). In order to consider how an electron transfer can take place we will study the following redox reaction:

2 Ag+ + Zn --> Zn2+ + 2 Ag

If a piece of Zn is placed in a solution of AgNO3, the reaction will occur spontaneously in which there is a direct transfer fo electrons from the Zn to the Ag+ (see Fig. 1). A small amount of work could be done, in principle, from the heat of the reaction.

If instead the Zn and Ag+ were separated as in shown in Figure 2, the same oxidation-reduction reaction would occur but with a flow of electrons through an external wire. The production of electrical or mechanical work is much more efficient for this process.

In Figure 2, the two half-cells are connected by a salt bridge, which prevents free mixing of both electrolyte solutions but permits the proper movement of ions to maintain electrical neutrality.

For example, when Zn is oxidized some anions must enter (or cations must leave) the Zn half-cell to compensate for the added positive charge of the Zn2+ produced. Also when Ag+ reduced cations must enter (or anions must leave) the Ag half-cell.

If oxidation-reduction reactions are split into half-reactions, it is possible to arbitrarily assign to each a voltage (the reduction potential) that indicates the relative tendency of that half-reaction to occur. The complete cell voltage (the only voltage that can be measured) is the difference between the two reduction potentials.

Referring to Figure 2, note that the anode half-reaction contains the better reducing agent (greater tendency to be oxidized) and therefore has the lower reduction potential. The anode is also the negative terminal in a galvanic cell. The anode is to be connected to the negative wire (black) from the voltmeter.

If oxidation-reduction reactions are split into half-reactions, it is possible to arbitrarily assign to each a voltage (the reduction potential) that indicates the relative tendency of that half-reaction to occur. The complete cell voltage (the only voltage that can be measured) is the difference between the two reduction potentials.

If oxidation-reduction reactions are split into half-reactions, it is possible to arbitrarily assign to each a voltage (the reduction potential) that indicates the relative tendency of that half-reaction to occur. The complete cell voltage (the only voltage that can be measured) is the difference between the two reduction potentials.

Referring to Figure 2, note that the anode half-reaction contains the better reducing agent (greater tendency to be oxidized) and therefore has the lower reduction potential. The anode is also the negative terminal in a galvanic cell. The anode is to be connected to the negative wire (black) from the voltmeter.

Procedure

Set the voltmeter to DC+ and 1.5 v (if your cell voltage is greater than 1.5 volts change to 5 v). Into separate beakers pour 0.10 M Cu(NO3)2, 0.10M Zn(NO3)2, 0.1 M FeSO4, and 0.10 M Pb(NO3)2, solutions to a depth of approximately 3 cm. Clean one strip each of Cu, Zn, Fe, and Pb using emery cloth or sandpaper. Place the metal strips into their respective solutions, i.e. the Cu goes into the Cu(NO3)2, the Zn into the Zn(NO3)2, etc. Carefully rinse a salt bridge (3 g of agar dissolved in 100 mL of 0.1 M KNO3) with deionized water and place it into the beakers containing the Zn and Cu. Connect the electrodes so that you get a positive voltage. Note the direction of electron flow (negative black to anode). Record the voltage and which electrode is the anode. Using the same procedure measure the voltage difference between the Zn half-cell and the Fe half-cell. The salt bridge should be carefully rinsed between measurements. Repeat for the Fe and Pb half-cells.

From the values you obtain in these three measurements, predict the voltage and direction of electron flow for the Pb|Pb(NO3)2||Cu(NO3)2|Cu cell. Measure the cell voltage and compare the experimental value with the predicted value.

Set up the Zn|Zn(NO3)2||Cu(NO3)2|Cu cell previously studied. Measure the voltage. Add, with stirring, 6 M NH3(aq) solution to the Cu(NO3)2 solution until the deep blue Cu(NH3)42+ complex is obtained. Read the voltage.

 

Results

CELL
VOLTAGE
ANODE

Zn|Zn(NO3)2||Cu(NO3)2|Cu

Zn|Zn(NO3)2||FeSO4|Fe

Fe|FeSO4||Pb(NO3)2|Pb

Pb|Pb(NO3)2||Cu(NO3)2|Cu

(predicted)

(measured)

Zn|Zn(NO3)2||Cu(NO3)2|Cu

 

above with NH3 added to

copper half-cell ________Explain the change in voltage

due to the additional ammonia.

Write the 4 reduction half-reactions in order of decreasing electrode potential (most positive on top and most negative on the bottom). If this is done correctly the best reducing agent will be on the bottom right and the best oxidizing agent on the upper left. If you assign a value of -0.76 volts to your Zn/Zn2+ electrode, calculate the electrode potentials of the other three.

Half-reaction (reduction)

Voltage

1

2

3

4

 

Les Brooks | Lecture Outline | Lab Outline | 115B Home Page

kh 8/9/99