CHM 130LL: Lab Organization and Policies

CHEM 162: Spontaneity as Viewed using Electrochemistry:

Galvanic, Concentration, and Electrolytic Cells

Introduction
A Galvanic Cell
The Nernst Equation
A Concentration Cell
An Electrolytic Cell
Procedure / In the thermodynamic study of the solubility of borax you determined that the free energy change for dissolving borax at 25 °C had a positive value. This result indicates that the process is non-spontaneous at 25 °C. Another interpretation is that a reactant-favored equilibrium between solid borax and its dissolved ions at 25 °C has the lowest free energy. Now, consider putting a battery into a flashlight and turning the light on. Is the free energy change from the battery’s point of view positive or negative? To answer this consider if the battery is doing work to light the bulb.
A spontaneous electrochemical cell (galvanic cell) is a type of battery that converts chemical energy into electrical energy. A galvanic cell has two connected half-cells in which each half-cell contains a metal electrode immersed in an aqueous solution of the same metal ion. The more active metal is the anode (where oxidation occurs, -) and the less active metal is the cathode (where reduction occurs, +). A wire, the external circuit, connects the two metal electrodes, and a salt bridge, the internal circuit, connects the solutions. The salt bridge is a partially contained aqueous solution of a non-reactive ionic compound. You will be constructing galvanic cells between Fe, Cu, Mg, and Zn half-cells.
When a galvanic cell operates, electrons, produced from the oxidized anode, spontaneously flow from the anode to the cathode via the external circuit. Dissolved cations collide with the cathode and are reduced by the transferred electrons. The cation concentration at the anode increases while it decreases at the cathode. The internal circuit or salt bridge corrects this imbalance; anions move toward the anode and cations move toward the cathode. As the cell operates the mass of the anode decreases while that of the cathode increases.

The symbolic representation of a galvanic cell constructed from Cr and Ni half-cells is: Cr(s) | Cr3+(aq) || Ni2+(aq) | Ni(s). By convention, the anode half-cell is on the left and is separated from the cathode half-cell on the right by the two parallel lines representing the salt bridge. A single vertical line separates different forms of the same substance. The sum of the half-reaction equations for oxidation and reduction yields the overall reaction equation and the number of electrons (n, moles) transferred.
2Cr(s)  2Cr3+(aq) + 6e- anode
3Ni2+(aq) + 6e-  3Ni(s) cathode

2Cr(s) + 3Ni2+(aq)  2Cr3+(aq) + 3Ni(s)
When the cell voltage (º) is positive the electrochemical reaction is spontaneous (Gº < 0,Gº = -nFº (F = 96,485 J/molV). A standard cell voltage (º), measured across the wire with a voltmeter, is the difference between the cathode’s reduction potential and the anode’s reduction potential (standard state conditions: 25 °C, 1 M, 1 bar). The more positive reduction potential applies to the cathode, where reduction occurs more favorably. Note that subtracting the anode reduction potential is the same as adding the oxidation potential, and is consistent with the oxidation occurring at the anode.
º = ºcathode –ºanode
Cr3+ + 3e-  Cr ºreduction = -0.73 V = ºanode
Ni2+ + 2e-  Ni ºreduction = -0.23 V = ºcathode
º = -0.23 –-0.73) = 0.50 V
Gº = -6 mol(96485 J/molV)0.50 V = 2.9 x 105 J
While Gº is proportional to the number of electrons transferred º is not. Remember that º is the potential that each electron experiences. An analogy is helpful here: one can roll 1 ball or 10 balls down a 10 m hill but each ball experiences the same change in height or potential energy.
The Nernst equation allows for measuring cell voltage under non-standard conditions. It is typically used when the initial concentrations of the ions are not 1 M. Note that Q is the initial mass balance relationship based on a balanced overall redox equation. In the standard Cr-Ni cell, = 1 and E= E.
where
A special type of a non-standard cell is a concentration cell. You will investigate the behavior of Cu-Cu concentration cell. A standard Cu-Cu cell potential would be zero. Diluting one of the half-cells creates a measurable potential difference. Changes to the concentration can also be made by complexing or precipitating the metal ions. You will observe how these concentration changes affect the voltage.
By entering E and E and the known concentration of one half-cell into the Nernst equation the unknown concentration of a test solution in the other half-cell can be calculated. This is the basis for determining pH using a pH meter and finding Ksp for nearly insoluble substances.
Your final investigation will involve the construction of an electrolytic cell. An electrolytic cell is a non-spontaneous cell that must have energy continuously inputted to force electrons from anode to cathode. Electrolysis is the method used to reduce ionic metals and oxidize ionic non-metals to pure elements. Electroplating a metal on another metal or plastic is one application of this method.
To convert the galvanic Cr-Ni cell into an electrolytic cell, the voltmeter is replaced by a power source that inputs more than 0.50 V to move electrons from the Ni to the Cr. This reversed flow makes Ni the anode (+) and Cr the cathode (-). Symbolically, changing a galvanic cell to an electrolytic cell simply swaps the names anode and cathode and reverses the flow of electrons.
You will be constructing a Cu-Cu electrolytic cell which reduces Cu2+ to Cu for the purpose of measuring Avogadro’s number (NA). After a period of reduction the mass gained by the Cu cathode will be measured and with this the moles of electrons used will be determined.

The total charge transferred will be calculated from the current (amperes = Coulombs/second) that flows for a measured time period (seconds) during the reduction. Given that the charge on 1 electron is 1.60 x 10-19 Coulombs, the total number of electrons transferred is:

Avogadro’s number is easily calculated from the quantity

Work in groups of 2, sign out the following equipment: a Fluke multimeter, a set of 20 mL beakers, a transformer, a U-tube, a LabQuest to use as a timer, and the necessary electrical wires with an alligator clip and a jack.
Galvanic Cells

Start the construction of the four different half-cells by filling each 20 mL beaker ¾ full with the appropriate 0.1 M aqueous solution; Cu(NO3)2, ZnNO3, MgSO4, and FeSO4 (made fresh). Complete the construction by: (1) sanding strips of copper, zinc, magnesium, and iron, (2) washing each strip with deionized water and then drying, and (3) folding the each strip over the lip of its respective beaker making sure the strip is partially immersed in the solution. These strips are the half-cell electrodes.

Prepare the 0.1 M FeSO4 solution by dissolving 0.54 grams of FeSO4 in 20 mL of DI water.

To prepare the Cu-Zn cell, clip one wire to each half-cell electrode. Attach the jack of the black wire to the COM port (black) and the red wire jack to the V port (red). This is the external circuit. Place a strip of filter paper on a watchglass and thoroughly moisten it with 0.1 M KNO3. Connect the half-cells by dipping the ends of the paper into the half-cell solutions. This is the internal circuit.

For galvanic cells the negative (-, black) electrode is the anode and the positive (+, red) electrode is the cathode. This set up will yield a positive voltage. If the voltage is negative, simply switch the jacks on the multimeter. The wire from the COM port connects to the anode. Set the multimeter to V direct current (V with the straight lines above it). Record your observations and voltage. Make sure to record the correct number of significant figures in the voltage.

Repeat the procedure to create a total of 6 cells. Make sure to use a new salt bridge for each cell.

A Concentration Cell
Do this portion of the experiment in the fume hood. All the reagents for this setup are in the fume hood.

Set up a galvanic cell using 1 M CuSO4 and 0.001 M CuSO4. Immerse sanded and washed Cu electrodes in each solution. Prepare a salt bridge and connect the two solutions. With the black wire in the COM port, connect the leads to the half-cells. After obtaining a positive voltage determine the anode and cathode and record the voltage. Respect significant figures.

Add 2-5 mL of 6 M NH3 to the 0.001 M CuSO4 solution until all precipitate dissolves (NH3 is an irritant, avoid inhalation). Record your visual and voltmeter observations. Respect significant figures.

Add 2-5 mL of 0.2 M Na2S to the 0.001 M CuSO4 solution, now containing NH3. Record your visual and voltmeter observations. Respect significant figures.

An Electrolytic Cell

Sand two Cu stripsand rinse with deionized water. Dry each piece thoroughly with a KimWipe and avoid touching with bare fingers. Label one of the pieces by scratching a mark near one of its ends. Measure the mass of each piece to the nearest 0.0001 g, respecting significant figures.

Dispense about 20 mL of 1.0 M CuSO4 (in 0.1 M H2SO4) into a 50 mL beaker. Slowly transfer this solution into the U-tube. Insert the copper electrodes into the U-tube and support them with the alligator clips of the leads. Make sure the marked electrode’s mark is above the solution. With the transformer unplugged from a power source set the voltage to 4.5 V and the polarity switch in the down position. Connect the Cu electrode with the smaller mass to the anode (+, red) of the transformer. Connect the transformer to the COM port of the multimeter using the black wire. Plug the red wire to the red “A” port on the multimeter. Connect the alligator clip from the red wire coming off the multimeter and attach to the cathode (-). Plug in the transformer and begin timing the electrolysis. Set the multimeter to the A (direct current). Do not move the electrodes during the electrolysis; this changes the current. Record the current at 1 minute intervals (an average current over the entire electrolysis period will be calculated).

After 15 minutes stop the electrolysis. Record the exact time (respect significant figures). Carefully remove the electrodes, particularly the plated Cu cathode. Rinse each electrode carefully by dipping it into a 400 mL beaker of deionized water then acetone. Air-dry each electrode and determine its mass to the nearest 0.0001 g. Record each mass respecting significant figures.

Return all borrowed equipment, chemicals, and metal strips.
*Discard remaining metal salt solutions in the labeled waste bottle in the waste hood.

Lab Report:
Complete the following in lab, obtain a staff signature, and submit.
1) Fill in the table with your observations.
Cell / E measured / Anode / Cathode / E° calculated
Cu-Zn
Cu-Mg
Cu-Fe
Zn-Mg
Fe-Mg
Zn-Mg
2) Write the balanced half-reaction equations and the balanced overall equation
for the Zn-Mg cell.
3) Give three reasons for the difference between E° calculated and E measured.
4) Fill in the table with your observations.
Cell / E measured / Observations / Anode / Cathode
initial
+ NH3 (aq)
+ Na2S (aq)
5) Write the balanced equation for the chemistry that occurs when NH3(aq) is added to one of the half-cells.
6) Explain the change in voltage going from the initial cell to the one containing NH3 (aq). Consider thermodynamics and equilibrium in your response.
7) Complete the table with your observations.
Quantity / Trial 1 / Trial 2
Initial Mass of Anode (g)
Initial Mass of Cathode (g)
Time of Electrolysis (s)
Average Current(A)
Final Mass of Anode (g)
Final Mass of Cathode (g)
Mass of Cu formed at Cat. (g)
Amount of Cu formed (mol)
Amount of e- transferred (mol)
Coulombs transferred (C)
e- transferred (#e-)
Calculated NA (#e-/mol e-)
Average NA
st. dev. in average NA
% relative uncertainty
% relative error
8) Attach your signed, photocopied lab notebook pages

7-Electrochemistry-v3-120208