2614 Electrochemistry and Galvanic Cells

1. INTRODUCTION

1.1 Objectives

After completing this experiment, the student will be able to:

• set up galvanic cells from several combinations of half-cells to determine anode, cathode, direction of electron flow, and electrochemical potential.
• vary concentrations of one solution in a half-cell to determine its effect on the electromotive potential.

1.2 Background

Spontaneous chemical reactions that involve the transfer of electrons can produce an electrical current in a galvanic cell. These oxidation-reduction reactions, called redox reactions, must physically separate the oxidation process of losing electrons from the reduction process of gaining electrons in two half-cells. The electrons generated then travel through an external wire from the electrode at the site of the oxidation, a piece of metal called the anode, to the electrode at the site of the reduction, another piece of metal called the cathode. To complete the electrical circuit, a salt bridge with a non-reacting ionic solution connects the two half-cells. Ions are able to flow out of the salt bridge to maintain electrical neutrality in each half-cell. See Figure 1 below.

Figure 1. A sample galvanic cell with Cu, Cu2+, Ag, and Ag+ reacting

Typically, the anode is displayed on the left and the cathode on the right, so the flow of electrons in the external wire is from left to right. In the figure, both metals are part of the overall reaction, but an inert electrode, usually a metal that is very non-reactive like platinum, may be used when the reactants are all dissolved species. The electromotive force, E, is measured with a voltmeter through which the electrons flow in the circuit.

The reaction at the anode is shown below.

Cu 🡪 Cu2+ + 2e-

The cathode reaction is shown below.

Ag+ + e- 🡪 Ag

The overall reaction is balanced, which means that the electrons must be equal on both sides and therefore cancel.

Cu + 2 Ag+ 🡪 Cu2+ + 2Ag

Because the anode is where electrons are generated, it is the negative electrode. Electrons are drawn to the cathode so it is the positive electrode. The electrons flow from anode to cathode spontaneously, but it is possible to connect the wires to the voltmeter incorrectly. A negative voltage simply means the wires have been attached incorrectly to the voltmeter.

While the electromotive force (emf) of a half-cell cannot be measured directly because it is an incomplete circuit, arbitrary values can be assigned to half-cells as long as the combinations which represent actual complete cells give the correct total electromotive force. Tables of these half-cell potentials are available. For the sake of consistency, these values are given for reduction potentials (Ered). The values listed are those for standard conditions: 298 K, 1 atm, and 1 M. Look up the standard half-cell potentials for the three half cells you will be studying and put their values on your data recording sheet (Table 1).

The standard electromotive force for a complete cell (Equation 1) is the difference in the reduction potentials for the half-cells.

E0cell = E0cathode – E0anode (Equation 1)

It is common to set up half-cells with concentrations that are not 1M. The value of E is then not the standard one but can be found from the Nernst equation (Equation 2).

Ecell = E0cell – RTnF ln Q (Equation 2)

n is the moles of electrons (electrons transferred in the reaction)

F is the Faraday’s constant = 96487 C/mol e-

R is the gas constant = 8.314 J/mol-K

T is the temperature in K

Q is the reaction quotient or equilibrium constant expression (but not necessarily the value of Keq)

In this experiment, several galvanic cells will be set up and their electromotive forces measured in volts. Also, the electromotive force of one cell with varying concentrations of the solution will be measured to confirm the Nernst equation.

2.0 SAFETY PRECAUTIONS AND WASTE DISPOSAL

3.0 CHEMICALS AND SolutionS

Chemical	Concentration	Approximate Amount	Notes
Zn, Cu, Fe metal	N/A	1 strip each	If they appear tarnished, clean with sandpaper until shiny
Zn(NO3)2	0.100 M	50 mL
Cu(NO3)2	0.100 M	50 mL
Fe(NH4)2(SO4)2 (ferrous ammonium sulfate)	0.100 M	50 mL	Keep bottle closed! Sensitive to air

4.0 GLASSWARE AND APPARATUS

100-mL beakers (3)	voltmeter
100-mL volumetric flask	ceramic cup
10.00-mL pipette	extra clean dry beakers

5.0 PROCEDURE

5.1 Measuring the Potential of a Galvanic Cell

1. Collect one strip each of Zn (dark silver metal), Cu (orange metal), and Fe metal (the Fe metal may be a nail). Make sure that these metal strips are clean by rubbing with sandpaper, if necessary.
2. Collect about 50 mL of each of the three metal solutions: Zn(NO3)2, Cu(NO3)2, and Fe(NH4)2(SO4)2 in clean and dry labelled 100-mL beakers.
3. You will not be using a salt bridge. You will be setting up your half cells in a slightly different manner (Figure 2). Pour a small amount of the Zn(NO3)2 solution inside the ceramic cup. Pour a small amount of the Cu(NO3)2 solution inside a 250 ml beaker (or other container of appropriate size). The solution level cannot be higher than the height of the ceramic cup. Place the ceramic cup inside the 250 ml beaker, taking care to keep the two solutions separated. The ceramic cup is porous, and allows for contact between the two solutions, but no appreciable mixing.

Figure 2. Set up for galvanic cell without a salt bridge

4. Using the red and black cables, attach the Zn and Cu strips to the voltmeter. Set the voltmeter to direct-current, and the voltage between 0 and 5 volts, or as appropriate to the meter. Dip the Zn electrode in the Zn(NO3)2 solution in the ceramic cup and the Cu electrode in the Cu(NO3)2 solution outside the cup. If the electrodes touch each other, a short circuit will occur, and there will be no voltage observed.
5. If the voltage reading is negative, you need to switch the clips between electrodes. (Because the voltmeter measures electrons flowing in a specific direction, one way shows a negative reading and the other positive. We know this is a spontaneous reaction however, so the Ecell is by definition going to be positive. Therefore, if we see a negative reading, we know we just need to switch the electrode clips to tell the voltmeter which is the correct cathode and which is the correct anode.) The reading may jump around a bit as you hold the clips, so take the highest reading and do it quickly. (The readings will decrease over time. Why?)
6. Record your readings on your data recording sheet (Table 2). Note which electrode is the negative one (anode) and record that also.
7. Clean the ceramic cell and beaker thoroughly with laboratory water and repeat steps 3 to 6 to set up the Zn-Fe cell combination and the Cu-Fe cell combination. It does not matter which solution is in the ceramic cup as long as the appropriate electrode is used with the appropriate solution.
8. Compare the theoretical values for the voltages to your experimental ones and give the percent errors on your data recording sheet (Table 2).

5.2 Nernst Equation and the Effect of Concentration on Ecell

1. Using your 10.00-mL pipette, take 10 mL of the Cu solution and place it in a 100-mL volumetric flask. Fill with laboratory water to the mark. Put on the stopper and thoroughly mix. Pour the solution into a clean dry beaker and label it as ‘Dilution 1’.
2. Thoroughly rinse the volumetric flask with laboratory water. It needs to be clean but not necessarily dry. Then, using a pipette, take 10-mL of the new diluted Cu solution (Dilution 1) and place them into the volumetric flask. Add laboratory water to the mark, mix well, and pour into a new clean and dry beaker labelled ‘Dilution 2’.
3. Repeat step 2 to make a third, fourth, and fifth dilution and label as Dilution 3, Dilution 4, and Dilution 5, respectively.
4. When all the diluted solutions are prepared, use each one of them with the Cu electrode as half-cell to connect with the Zn half-cell. Set up the galvanic cell as before (Part 5.1). Be sure to clean the ceramic cup thoroughly between measurements. Record the voltage you measure for each combination of the original Zn half-cell with Dilutions 1 through 5 of the Cu(NO3)2 solution on your data recording sheet (Table 3).
5. On Table 3, record the concentration of the Cu2+ ions in each of the dilutions and the predicted value of the voltage when combined with the Zn half-cell according to the Nernst equation (Equation 2). Compare the results of experimental and theoretical values and determine the percent error. (Remember % error = |(observed value-theoretical)/(theoretical)| * 100%)

6.0 DATA RECORDING SHEET AND CALCULATIONS

Last Name

First Name

Partner Name(s)

Date

Table 1. Standard Value of Electromotive Forces

Half-cell Reaction	E0red
Cu2+ + 2e- → Cu
Fe2+ + 2e- → Fe
Zn2+ + 2e- → Zn

What source did you use for these values? _______________________________

Table 2. Experimental and Theoretical Values of Eocell

Cell	Which electrode is the anode?	Experimental Eocell	Theoretical Eocell (Eocathode – Eoanode)	% Error
Zn-Cu
Zn-Fe
Cu-Fe

Show calculations for one cell below.

Table 3: Effect of Concentration on Ecell

Dilution	[Cu+2]	Experimental Ecell	Theoretical Ecell	% Error
1
2
3
4
5

Show calculations for one cell below.

7.0 POST-LAB QUESTIONS

1. It was indicated in the procedure that the voltage would drop over time, eventually reaching zero. Why?

2. All solutions used in this experiment were 0.100 M, which is not standard. Did that cause our experimental values to be different from the theoretical E0cell? Why or why not?

3. Could have any other factor caused our experimental values to be different from E0cell? Explain your answer briefly.

4. Based on your data in Part 6.0, what trend did you observe on the effect of concentration on the value of Ecell?