Galvanic Cells Lab Report

Christa Schafer
[e-mail]
Chemistry 125.339
College of LSA 2018
University of Michigan

Galvanic Cells
Abstract:

The purpose of this lab was to determine the concentration of an unknown copper solution using galvanic cells and the flow of electrons from chemical energy into electrical energy. Our hypothesis was that using the oxidation/reductions reaction in a galvanic cell that occur from the transfer of electrons, we would be able to determine the concentration of an unknown copper solution. In order to do this, Lab 10 was broken up into three sub-labs referred to as Lab 10A, Lab 10B, and Lab 10C, respectively. In Lab 10A, the objective was to determine the reduction potential for iron. This was done by submerging different

metals into their respective ion solutions and connecting two of these solutions using a galvonic cell to determine reduction potential. Through calculation, the reduction potentail was determined to be 467 mV in the experiement with zinc and 511 mV in the experiement with copper. In Lab 10B, the objective was to calculate the voltage of the reaction between two different molar solutions of copper (II) using the Nernst equasion.This was preformed by establishing a galvanic cell using copper (II) in two different molarities.The result of voltage, 80 mV was compared to the actual reduction potential for copper (II). In Lab10C, the objectvie was to determine the concentration of the unknown.This was achieved by measuring the half cell potential for different concentration of the unknown copper and zinc solid. The resultign concentration was determined to be .23 M.

Introduction:

Electricity is the movement of electrons from one point to another1. When this flow of electrons is spontaneous between two solutions, it is referred to as a galvanic cell1. The energy of a galvanic cell is harnessed when chemical energy is converted to chemical energy in the transfer of electrons. This energy from the transfer of electrons is the same energy that is used to power an iPod or toy battery. The cell is constructed by connecting the positive end of a multimeter with the cathode and connecting the negative end of the voltmeter to the anode1. The electrode gaining the electrons, therefore being reduced is the cathode. While the electrode that is providing the electrons is the anion, the electrode that is bring reduced. In order for the reaction to proceed, there also needs to be the presence of a salt bridge to provide cations to the cathode (which has an excess amount of electrons) and provide anions to the anode (which has an excess amount of electrons). The potential of a reaction to occur at a standard state is known as the standard reduction potential2. The quantitative values for the potentials are summarized in the table for standard reduction potential. The more negative a standard reduction potential value is, the less likely it is for a reaction to occur spontaneously. However, the more positive the standard reduction potential value is, the more likely a reaction will occur. These values can be used to compare and identify the likelihood of reaction and voltage that results.

Experimental:

In the initial part of the experiment, Lab 10A, in which the reduction potential for irons was being identified, begin by sanding two strips of copper, one strip of zinc and one strip of iron to act as your electrodes. Fill one 50-ml. beaker to 30 ml. with 0.1 M Zn(NO3)2, a second 30 ml full with 0.1 M CuSO4 and a third 30 ml. full with 0.1 M FeSO4. Place the one of the copper strips into in the beaker with 0.1 M CuSO4. Similarly, place the zinc strip into the beaker with 0.1 M Zn (NO3) 2. Roll a large piece of filter paper and saturated with 0.1 M KNO3 solution and fold so that each end of the filter paper is in one of the solutions. Connect two leads to the multimeter, then the probe end of each lead to the electrodes so that the red probe is attached to the copper solution. Set to a range of 2000 mV. Record measurement of voltage to be 806 mV for the copper and zinc solution, 293 mV for the copper and iron solution, and 851 mV for the zinc and iron solution. Repeat using every combination of the three solutions, ensuring to prepare a new salt bridge each time.

A Galvanic Cell using Copper (II) and Zinc (II) 2

In the second part of the lab, Lab 10B, a galvanic cell using 1.0 M CuSO4 and 0.001 M CuSO4 is established and a polished copper electrode was placed in each. Attach the red probe to the 1.0 solution of CuSO4 and measure the voltage to receive a result. The experimental voltage was determined to 88mV.

In Lab 10C, set up four more galvanic cells, this time using the Zn/Zn2+ half-cell each time, and pairing it with the 0.1 M. 0.01 M, 0.001 M, and 0.0001 M CuSO4 half-cells in sequence. Connect two leads to the multimeter, then the probe end of each lead to the electrodes so that the red probe is attached to the copper solution. Set to a range of 2000 mV, or 20 V. Record measurement of voltage. Repeat using every combination of molarity, ensuring to prepare a new salt bridge each time. In the 0.1M solutions, the potential was measured to be 0. 59 V, in the .01 solution it was measured to be 0.58V, in the .001 it was measure to be 0.54 V and in the .0001 M solution it was measured to be 0.50 V.

Data and Calculations:
Lab 10A:

Reduction Potential Measurements for Transition Metals*

Substances Multimeter Measurement (mV)
0.1 M Zn/Zn 2+ and 0.1 M Cu+2/Cu 806
0.1 M Zn/Zn 2+ and 0.1 M Fe+2/Fe 293
0.1 M Fe+2/Fe and 0.1 M Cu+2/Cu 851

*Due to inoperable lab error, lab data used for Lab 10A from another group

Reaction of Zinc and Copper

Oxidation Half Reaction:
Zn → Zn2+ + 2e- EO =-760 mV

Reduction Half Reaction:
Cu+2 + 2e-→ Cu EO =340 mV

Theoretical Ecell Yield:
340 mV- (-760 mV) = 1.1 V

Actual Ecell Yield:
EO =806 mV

Percent Error:
(806mV-1100mV)/1100mV=26.7% Reaction of Zinc and Iron

Oxidation of Half Reaction
Zn → Zn2+ + 2e- EO =-760 mV

Reduction Half Reaction:
Fe+2 + 2e-→ Fe EO = X

Reduction Potential for Iron with Zinc:
293 mV-760 mV = -467

Percent Error:
(467mV-440)/440mV=6.1% Reaction of Copper and Iron

Oxidation of Half Reaction
Fe → Fe+2 + 2e- EO = X

Reduction Half Reaction:
Cu+2 + 2e-→ Cu EO =340 mV

Reduction Potential for Iron with Zinc:
851 mV – 340mV= 511 mV

Percent Error:
(511mV-440)/440mV=16%

Lab 10B:
Reduction Potential Measurement for Copper Sulfate of Varying Molarity
Substance Multimeter Measurement (mV)
0.001 M Cu+2/Cu and 1.0 M Cu+2/Cu 80

Oxidation Reaction:

(0.001 M) Cu → Cu2+ + 2e –

Reduction Reaction

(1M) Cu+2 + 2e-→ Cu

Percent Error:
(88mV-80mV)/88mV=9.1%

Nernst Equation:
Ecell= EO cell – (.0592/n) x log (Q)
Q = log (.001M/1.0M)= -3
N=2

Theoretical Ecell= 0
Ecell= 88 mV

Lab 10C:
Reduction Potential Measurements of Copper and Zinc
Concentration of CuSO4 Multimeter Measurement (mV) -Log(Concentration)
0.1 590 1
0.01 580 2
0.001 540 3
0.0001 500

4

Reduction Potential Measurement for Unknown Copper B
Substance Multimeter Measurement (mV)
Unknown Copper B 610

Equation of Line (Graphing Calculator)

Ecell= -31 (-log (concentration))+630
610=-31 (-log (concentration))+630
.23 M = concentration

Results:

The purpose of this lab was to determine the concentration of an unknown copper solution using galvanic cells and the flow of electrons from chemical energy into electrical energy. We were able to determine that the unknown copper solution was 0.23 M.

In Lab 10A, the objective was to determine the reduction potential for iron. This was done by submerging different metals into their respective ion solutions and connecting two of these solutions using a galvonic cell to determine reduction potential. Due to a complication in the lab, assumably inproper calibration of the multimeter, our lab results needed to be substittued by another group’s data for Lab 10A only.Through calculating the Ecell for mutiple oxidation/ reduction reactions and using the standard reduction potential we aleady know we were able to determine that the reduction potentail was 467 mV in the experiement with zinc and 511 in the experiement with copper. The two resulted in 6.1% error and 16% error, respectivly. Some of our error could have been the result of using another group’s data who did not use the same mutlimeter we used, as welll as from the room not being in standard state. The values we collected for standard reduction potenetial display ideal conditions in a 25 degree C environment, which we did not have.

In Lab 10B, the objective was to calculate the voltage of the reaction between two different molar solutions of copper (II) using the Nernst equasion.This was preformed by establishing a galovnic cell using copper (II) in two different molarities.The resulting voltage, 80 mV was compared to the actual reduction potential for copper (II). The acutal voltage, 88 mV, proved to only be slightly different from our calculation, producing a perecent error of 8%. This percent error could be the result of not being in a standard environment or of error in the innaccuracry of the mutilmeter.

In Lab10C, the objectvie was to determining the concentration of the unknown.This was achieved by measuring the half cell potential for different concentration of the unknown copper and zinc solid. From the voltages that were producted from the cell, an equasion was derrieved using the Nernst equation from the realationship bewteen the negatvie log of the concentation and the voltage ( in mV). The resultign concentration was determined to be .23 M from the formula and the uknown’s voltage. There was no conlcusion that was able to be made on the accuracy of the calculated molarity. However, any error that does exist could have arose from the amibigity and uncertainity of the calculated values and from the inacccuracy of our dilutions.

Conclusion:
Overall, we were able to meet the objective for the lab by calculating the concentration of the unknown copper solution. We were unable to draw conclusions regarding the accuracy of our resulting concentration because we did not have the true value for the concentration. As labeled above, the solutions that lost electrons were oxidized, and therefore were the anodes. The cathodes were the solutions that were reduced and gained electrons. When the anode loses electrons they are transferred to the cathode through the metal wires connecting the two solutions and multimeter. In Lab 10A, the electrons traveled from the zinc to the copper (II), from the zinc to the iron (II) and the iron (II) to the copper (II) in their respective reactions. In Lab 10B, the electrons progressed from the 0.0001 M copper solution to the 1.0 M solution. Lab 10C resulted in electrons being transferred from the zinc to the copper in all reactions. For each reaction in all labs, the salt bridge provided stability for the solutions when the electron count progressed too far. In all reactions, the salt bridge provides cations to the cathode in order to balance out the build up of electrons from the electron flow. The anions move to balance of the charge of the anode in all reactions. From the Nernst equation we are able to see that, a difference in concentration can completely alter the expected Ecell value. The molarity with the smallest/ most negative Ecell, or lowest molarity will be likely oxidized, while the other acts as the oxidizing agent and allowing for a oxidation reduction reaction to still occur.
Acknowledgments:
I would like to send my deepest thanks to my lab partner, Vicki Kazdaglis, as well as to my professor, Dr. Alexander Poniatowski, and my Graduate Student Instructor, Rosina Ho Wu.
References:
1."Standard Reduction Potential." - Chemwiki. University of California, Davis, n.d. Web. 20 Nov. 2014.

2. Jespersen, Neil D., and Neil D. Jespersen. "Oxidation- Reduction Reactions and Electrochemistry." AP Chemistry. Hauppauge, NY: Barron's Educational Series, 2010. N. pag. Print.

3. " Standard Reduction Potentials at 25 C." STANDARD REDUCTION POTENTIALS IN AQUEOUS SOLUTION AT 25 ° C (n.d.): n. pag. AP Central. College Board. Web.