Electrochemistry

Introduction

The course description for the College Board AP^® Chemistry lists the following topics for electrochemistry; electrolytic and galvanic cells, Faraday’s laws, standard half-cell potentials, Nernst equation, and prediction of the direction of redox reactions. Use this set of three electrochemistry demonstrations to both engage your students in a hands on review and test their understanding of this sometimes difficult yet fundamental topic.
The set of three demonstrations includes:

1. The Voltaic Cell—Begin with the “textbook” zinc—copper redox reaction. In this demonstration, students will use the table of standard electrode potentials and the Nernst equation to correctly assess direction of the current and then the change in current when the electrolyte concentration is reduced.
2. Electrolysis—By observing the changes in pH, students will use Faraday’s law to determined the average current of an electrolysis reaction.
3. Concentration Cell—Nothing challenges understanding of the Nernst equation like the concentration cell.

Experiment Overview

The Voltaic Cell
The voltaic cell is an electrochemical cell in which an electrical current is generated from an oxidation—reduction reaction involving electron transfer. In this demonstration, challenge your students’ understanding of this fundamental topic of electrochemistry.

Electrolysis
Reactant-favored oxidation reduction reactions can be made product-favored by the use of electrical energy. Set up an electrolysis reaction with a 6-volt battery and a solution of potassium iodide. After 15 minutes, take the pH reading of the solution. From the knowledge of starting solution, their observations of the reactions, and the final pH, students will determine the average current used during the electrolysis.

Concentration Cell
A voltaic cell can be constructed using identical half cells where ε°_Cell equals zero. All that’s needed is a difference in concentration of the metal ion in the solutions.

Materials

Safety Precautions

Potassium nitrate in solid form is a strong oxidant and a fire and explosion risk when heated or in contact with organic material. It is also a skin irritant. Zinc sulfate solution is a skin and eye irritation. Copper(II) sulfate solution is a skin and respiratory irritant; it is slightly toxic by ingestion. The electrolysis reaction generates a small amount of iodine gas. Perform this demonstration in a well-ventilated lab only and do not breathe the vapor. Avoid contact of all chemicals with skin and eyes. Zinc sulfate solution is a skin and eye irritation. Wear chemical splash goggles, chemical-resistant gloves and a chemical-resistant apron. Please review current Safety Data Sheets for additional safety, handling and disposal information.

Disposal

Please consult your current Flinn Scientific Catalog/Reference Manual for general guidelines and specific procedures, and review all federal, state and local regulations that may apply, before proceeding. The copper and zinc sulfate solutions may be reused or flushed down the drain with excess water according to Flinn Suggested Disposal Method #26b. The electrolysis reaction generates an iodine-water solution. This solution should be reduced with sodium thiosulfate according to Flinn Suggested Disposal Method #12a. The resulting waste solution should be allowed to sit overnight to thoroughly degas. It may then be rinsed down the drain with plenty of excess water according to Flinn Suggested Disposal Method #26b. Do not dispose of the waste solution directly down the drain. The zinc sulfate solutions may be reused or flushed down the drain with excess water according to Flinn Suggested Disposal Method #26b.

Prelab Preparation

The Voltaic Cell

Salt Bridge

1. Add 150 mL of distilled or deionized water to a 400-mL beaker and place it on a hot plate.
2. Heat the water to boiling, and then add 3 g of agar. Stir the mixture until it forms a uniform suspension.
3. Use beaker tongs to remove the beaker from the heat. Add 15 g of potassium nitrate to the beaker and stir until the KNO₃ dissolves.
4. Invert the U-tube and carefully fill the tube with the KNO₃/agar mixture.
5. Use a ring stand and buret clamp to hold the U-tube vertically and let the tube set up overnight.

Procedure

The Voltaic Cell

1. Set up the cell in Figure 1.
   {12810_Procedure_Figure_1}
2. Add 200 mL of 1 M zinc sulfate solution a clean 400-mL beaker and 200 mL of 1 M copper sulfate to another clean 400-mL beaker.
3. Place the salt bridge in both beakers (see Figure 1).
4. Dip both the copper and zinc strips in a solution of 6 M nitric acid, then rinse the strips with distilled or deionized water.
5. Connect the zinc electrode to one end of an alligator clip; do the same with the copper strip and another alligator clip.
6. Before connecting the voltmeter, ask the students to write the Nernst equation for this voltaic cell and determine which electrode should be the cathode and which should be the anode and what voltage should they see when the voltmeter is connected.
7. Connect the zinc to the negative lead (anode) and the copper strip to the positive lead (cathode). Read the voltage.
8. Disconnect the leads from the voltmeter.
9. Now pour approximately 200 mL of the 0.01 M zinc sulfate into a clean 400-mL beaker.
10. Remove the 1 M zinc sulfate beaker from the setup. Rinse off the zinc electrode with deionized water.
11. Tell the class that the 1.0 M zinc sulfate solution is being replaced by a 0.01 M zinc sulfate solution. Before connecting the new cell to the voltmeter, instruct the students to first predict whether the voltage will increase, decrease or remain unchanged, and then, if the voltage does change, calculated the magnitude of that change.
12. Construct the voltaic cell as in Figure 1, using the beaker containing the 0.01 M zinc sulfate solution. Have the students record the new voltage.

1. Place the 600-mL beaker containing the KI solution on the ring stand.
2. Use the clamps to suspend the two carbon electrodes in the potassium iodide solution. Make sure the electrodes are not touching (see Figure 2).
   {12810_Procedure_Figure_2}
3. Connect one end of each the alligator clip lead to a separate carbon electrode. Connect one alligator lead to the positive post of the battery.
4. Use a pH meter to take the initial pH of the solution. Have students record this value.
5. Connect the second lead to the negative post of the battery. Have students record the time.
6. Have students make observations of the reaction. Note any color changes and any gas formation.
7. Let this run for another 15 minutes. Proceed to the next demonstration while the electrolysis continues.
8. After 15 minutes, disconnect a wire to the battery and have the students record the time.
9. Take a pH reading. Have students record the value.

1. Set up the cell in Figure 3.
   {12810_Procedure_Figure_3}
2. Add 200 mL of 0.01 M zinc sulfate solution a clean 400-mL beaker and 200 mL of 1 M zinc sulfate to another clean 400- mL beaker.
3. Place the salt bridge from the voltaic cell demonstration in both beakers (see Figure 3).
4. Dip both the zinc strips in a solution of 6 M nitric acid, then rinse the strips with distilled or deionized water.
5. Connect a zinc electrode to one end of an alligator clip leads; do the same with the other zinc strip and another alligator clip.
6. Before connecting the voltmeter, ask the students to write the Nernst equation for this voltaic cell and determine which electrode should be the cathode and which should be the anode and what voltage should they see when the voltmeter is connected.
7. Connect the zinc strip in the more concentrated solution to the negative lead (anode) and the other copper strip to the positive lead (cathode). Read the voltage.

Student Worksheet PDF

12810_Student1.pdf

Teacher Tips

• When using a voltmeter in the Voltaic Cell demonstration, make sure the connections are tight and the metal electrodes are shiny. The maximum voltage will be less than 3 volts. Use the smallest range that gives a reading that is on the scale. Be sure to use the connections for DC voltage. If the voltmeter gives a negative voltage, reverse the connections. The cathode connects to the positive terminal, the anode to the negative terminal.
• The most convenient type of voltmeter to use is a digital voltmeter, but an analog meter may also be used. Borrow voltmeters from the physics department, if necessary. Review the specific operation of the voltmeter before beginning the demonstrations.
• Students may need help in understanding the calculations using the Nernst equation. Remember that solid substances do not appear in the reaction quotient expression, Q, and only ion concentrations are in the values of Q.
• Many different questions can be posed to test student understanding of the electrochemistry of the voltaic cell.
  1. What would happen to the voltage reading if a palladium wire was substituted for the salt bridge?
  2. Will the voltage increase, decrease or stay the same if...

     50 mL of 2 M Cu(II) SO₄ solution were added to the right beaker?
     A large bar of copper were substituted for the copper strip?

• You may want to take the students through the Electrolysis calculations step by step. Make sure they understand that pOH represents 10^–[OH] (ten to the minus hydroxide ion concentration, in moles per liter).
• Variations of this demonstration include other electrolytes, such as silver nitrate and zinc bromide. Both these metals are more easily reduced than hydrogen atoms in water. Electrolysis of silver nitrate generates silver metal at the cathode. Zinc bromide generates zinc metal at the cathode and bromine at the anode. In both cases, the deposited mass of metal will serve as the appropriate data for calculating coulombs.
• Potassium iodide solution is light- and air-sensitive. Prepare the solution fresh within two weeks of its anticipated use and store the solution in a dark bottle, if possible.
• A variation on the Concentration Cell demonstration uses copper metal and solutions of copper(II) sulfate that have identical concentrations. Sodium sulfide, Na₂S, is added to one beaker of copper sulfate. This results in the precipitation of copper sulfide, CuS. Cu(II) ion concentration is thus reduced, and a copper ion concentration cell is set up.

Sample Data

Electrolysis 

Data Table 
Initial pH ___6.5___ 
Elapsed time: ___14___ min. ___26___sec. 
Final pH ___10.6___ 

Answers to Questions

The Voltaic Cell

1. Using the table of standard reduction values, determine ε°_cell for this voltaic cell arrangement?

   The standard reduction potentials for copper(II) and zinc ions are:
   Cu²⁺(aq) + 2e^– → Cu(s)         Cu²⁺(aq) + 2e^– → Cu(s)          ε° = 0.337 V
   Zn²⁺(aq) + 2e^– → Zn(s)             ε° = –0.763 V
   
   Since copper has a higher value of ε°, the voltaic cell will have copper(II) ions reduced at the cathode and zinc metal oxidized at the anode. ε° for the cell will be
   ε°_cell = ε°_cathode – ε°_anode
   ε°_cell = 0.337 V – (–0.763 V)
   ε°_cell = 1.10 V

2. Write the Nernst equation for this cell.
   {12810_Answers_Equation_1}

   If the overall reaction for the voltaic cell is
   Cu²⁺(aq) + Zn(s) → Cu(s) + Zn²⁺(aq)
   then Q = [Zn²⁺]/[Cu²⁺] and

   {12810_Answers_Equation_2}
3. Before the leads are connected, predict the results for the new voltage value when a 0.01 M zinc sulfate solution is substituted for the 1 M zinc sulfate solution.

   If the zinc concentration decreases, the Q value becomes smaller and less than one. The natural log of any number less than one is a negative number. This makes the expression

   {12810_Answers_Equation_3}
4. With the right beaker containing 1 M zinc sulfate solution, what would happen to the voltage if 50 mL of distilled water were added to the left beaker containing 1 M copper sulfate solution?

   Adding water decreases the copper ion concentration, making the Q value becomes larger and greater than one. The natural log of any number greater than one is a positive number. This makes the expression

   {12810_Answers_Equation_4}

1. For this electrolysis reaction, the species present in solution are potassium ions, K⁺, iodide ions, I^–, and water molecules. The standard reduction or oxidation potential for these species are:
   {12810_Answers_Table_1}
   Based on these oxidation and reduction potential values, write the half cell reactions and the overall oxidation–reduction reactions that occur.

   Combine the highest reduction potential with the highest oxidation potential.

   {12810_Answers_Table_2}
2. Recall that current is the flow of electrons and is measured in amperes, A. One ampere represents the flow of one coulomb of charge per second. Faraday’s constant relates the coulomb to electrons and has the value of 96, 500 coulomb/mole e^–.

   Refer to the oxidation–reduction equation for this electrolysis. Relate the number of moles of electrons, e^–, transferred to the number of moles of hydroxide ions, OH^–, produced.
   The number of moles of electrons transferred is equal to the moles of hydroxide ions produced.

3. Based on the pH values, calculate the initial and final hydroxide ion concentrations, [OH^–].

   [OH^–] = 10^–pOH ; pOH = 14 – pH

   The initial pH is 6.5, the initial pOH is 14 – 6.5 or 7.5
   The initial [OH^–] is equal to 10^–7.5 or 3.2 x 10^–8 moles/L
   
   The final pH is 10.6, the final pOH is 14 – 10.6 or 3.4
   The final [OH^–] is equal to 10^–3.4 or 4.0 x 10^–4 moles/L
   
   How many moles of hydroxide ions were produced? How many coulombs of electrons were transferred?
   Since the final concentration of hydroxide ions is very much greater than the initial concentration, it can be assumed that

   [OH^–]_Final – [OH^–]_Initial = [OH^–]_Final

   The moles of hydroxide ions produced is equal to the concentration times the solution volume.

   { (4.0 x 10^–4) moles/L} x 0.2 L = 8.0 x 10^–5 moles OH^–

   Moles of electrons transferred is equal to the moles of hydroxide ions produced.

   What was the average current, A, for this electrolysis?
   A = coulombs/total elapsed time (sec)
   Total elapsed time = 14 min. 26 sec.

   = {(14)(60) + 26}sec
   = 866 sec

   {12810_Answers_Equation_5}

   Coulombs = 7.8 coulombs
   A = 7.8 coulombs/866 sec = 0.009A

1. What is the ε°_cell value for this voltaic cell arrangement?

   Since both the anode and cathode are zinc/zinc ion half cells, ε°_cell is zero.

2. Write the Nernst equation for this cell.
   {12810_Answers_Equation_6}
3. Is this cell spontaneous? If so, where is the anode? Cathode?

   For the cell to be spontaneous, ε_cell must be positive.

   {12810_Answers_Equation_7}

   ε_cell = – 0.0124 ln(0.01) = +0.057V

   The cathode is the cell with the zinc ion concentration of 1 molar.

   The anode is the cell with the zinc ion concentration of 0.01 molar.

4. What would happen to the voltage if 50 mL of distilled water were added to the left beaker containing 0.01 M zinc sulfate solution?

   A lower concentration of zinc ions at the would decrease the Q value([Zn²⁺]_A/[Zn²⁺]_C). As Q decreases, the expression –lnQ increases or ε_cell increases.

Discussion