This presentation illustrates how chemical reactions can generate electricity. This presentation uses a series of real concentration electrochemical cell and a series of concentration cells from a computer simulation (AACT) to help students understand the concepts associated with non-standard electrochemical cells and the use of the Nernst equation. Electrochemical cells based on an electromotive force (emf) generated by virtue of a difference in concentration between two half-cells having different types of metal electrodes and corresponding aqueous solutions not at 1.0 M, called a Nernst Cell. These are electrochemical cells that operate with a difference in concentration between two half-cells having the different types of metal electrodes, and the solutions are not at standard conditions (i.e. not 1.0 M).
A zinc-copper voltaic cell with non-standard concentrations of aqueous solutions yields a cell voltage less than +1.10 volts, +0.89 Volts. In this cell the concentration of zinc ions in solution in the beaker on the left is greater than the concentration of Cu2+ ions in solution in the beaker on the right. [image from MrGrodskiChemistry, accessed, May 2023]
A second type of non-standard electrochemical cell has the same type of metal electrode in corresponding metal ion aqueous solutions not at 1.0 M. This type of cell is called a concentration cell. For example, a half-cell of dilute solution of copper(II) sulfate with a copper electrode is connected by a salt bridge to a half-cell with concentrated copper(II) sulfate with a copper electrode. The emf of the cell is, E = +0.58 Volts.
Image from John Straub's Lecture notes, Boston University.
https://people.bu.edu/straub/courses/demomaster/concentrationcell.html [accessed May, 2023]
Summary - what this presentation offers with respect to learning electrochemistry
- Electrical energy can be generated by chemical reactions occuring due to a difference in concentration of solutions in half-cells.
- Connects among the three levels of representation: macroscopic (physical electrochemical cells), symbolic (chemical equations, cell diagrams), and the microscopic - particle level (particle level drawings, animations at the particle level).
- Introduction and application of the Nernst equation. Understanding the Nernst equation at a conceptual level.
This presentation also provides an opportunity for students to make connections among the macroscopic, particle level, and symbolic levels of representation (Johnstone, 1982, 1991, 1993) associated with electrochemical cells.
This presentation/lesson is consistent with the principles of Universal Design for learning in that multiple representations are incorporated in the presentation. This presentation also provides an opportunity for students to make connections among the macroscopic, particle level, and symbolic levels of representation (Johnstone, 1982, 1991, 1993) associated with electrochemical cell processes. The AACT simulation has particle level animations of what occurs at the surface of the electrodes, migration of ions in the aqueous solution, oxidation and reduction half-reactions, and migration of ions in the salt bridge. Also, identification of which electrode is connected to the ground terminal and which electrode is connected to the "lightning bolt" terminal of the voltmeter.
Nernst Cells AACT Simulation
Metals: copper, zinc, magnesium, silver.
Aqueous solutions: copper(II) nitrate, zinc nitrate, magnesium nitrate, silver nitrate concentrations: 2.0 M, 1.0 M, 0.10 M, 0.010 M, 0.0010 M
For example, zinc metal electrode in 2.0 M Zn(NO3)2 solution, a copper metal electrode in a 0.0010 M Cu(NO3)2 solution, and a connecting salt bridge. The electrodes are connected to a voltmeter. E cell = +1.00 Volts.
https://teachchemistry.org/classroom-resources/galvanic-voltaic-cells-2 [accessed May, 2023]
ACS AACT computer simulation. Greenbowe, T.J.; Gelder, J.I., Boyd, A, Wixon, M. (2020). Galvanic/Voltaic Cells 2. American Association of Chemistry Teachers, American Chemical Society: Washington, D.C.
The following diagram identifies the components of a zinc-cooper Nernst electrochemical cell.
This presentation also provides an opportunity for students to make connections among the macroscopic, particle level, and symbolic levels of representation (Johnstone, 1982, 1991, 1993) associated with the dissolving process.
This presentation/lesson is consistent with the principles of Universal Design for learning in that multiple representations are incorporated in the presentation. This presentation also provides an opportunity for students to make connections among the macroscopic, particle level, and symbolic levels of representation (Johnstone, 1982, 1991, 1993) associated with electrochemical cell processes. The AACT simulation has particle level animations of what occurs at the surface of the electrodes, migration of ions in the aqueous solution, oxidation and reduction half-reactions, and migration of ions in the salt bridge.
https://teachchemistry.org/classroom-resources/galvanic-voltaic-cells-2 [accessed May, 2023]
ACS AACT computer simulation. Greenbowe, T.J.; Gelder, J.I., Boyd, A, Wixon, M. (2020). Galvanic/Voltaic Cells 2. American Association of Chemistry Teachers, American Chemical Society: Washington, D.C.
Students use an activity sheet to write net ionic equations, indicate the electron flow in the wires and metal electrode and indicated the direction of migration of cations and anions within the salt bridge.
An activity sheet for this instructional event is available from the AACT web site.
Web page author: T. Greenbowe, University of Oregon. This page is under construction.
Concentration Cells
The second type of non-standard electrochemical cell has the same type of metal electrode in corresponding aqueous solutions not at 1.0 M. This type of cell is called a concentration cell. In the image below, a half-cell of dilute solution of copper(II) sulfate with a copper electrode is connected by a salt bridge to a half-cell with concentrated copper(II) sulfate with a copper electrode. The emf of the cell is, E = +0.58 Volts.
Image from John Straub's Lecture notes, Boston University.
https://people.bu.edu/straub/courses/demomaster/concentrationcell.html [accessed May, 2023]
The following diagram illustrates the components of a copper concentration cell. One half-cell has a copper electrode in 2.0 M copper(II) sulfate solution. The other half cell has a copper electrode in 0.0010 M copper(II) sulfate solution. A salt bridge connects the two half-cells. A voltmeter measures the potential difference between the two half cells, +0.10 Volts.
In the Galvanic/Voltaic Cells 2 computer simulation posted on the ACS AACT web site, one can construct a variety of non-standard voltaic cells. For each half-reaction, choose a metal and select the corresponding metal ion concentration of solution. Record the cell potential, E, from the voltmeter. Animations representing the processes that occur at the particle level at each electrode, and in the salt bridge. In the following copper concentration cell diagram, one half-cell has an initial concentration of 0.001 M Cu2+ and the other has a concentration of 2.00 M Cu2+. The emf of the cell is +0.10 Volts.
After some time, the two half-cells reach the same concentration of Cu2+ ions and the voltage is 0.00 Volts.
ACS AACT Computer Simulation: Galvanic/Voltaic Cells 2. Greenbowe, T.J.; Gelder, J.I., Boyd, A, Wixon, M. (2020). Galvanic/Voltaic Cells 2. American Association of Chemistry Teachers, American Chemical Society: Washington, D.C.
Ask students the following questions: "If we want to maximize the voltage from this type of cell, how should we select the concentration of ions?" "In which directions do electrons flow in this cell?" "What half-reactions occur at each electrode?"
Nature has a tendency to attempt to make the concentrations of the connected solutions to be equal. As the reaction proceeds, we expect the dilute solution to become a bit more concentrated and the concentrated solution to become a bit dilute. Given enough time, the two solutions should have equal concentrations.
The way for this to occur is for copper metal atom from the electrode in the dilute solution to be reduced to copper(II) ions, thus increasing the Cu2+ ion concentration in solution. The copper electrode decreases in mass.
The reaction occuring at the anode will be an oxidation half-reaction
Cu(s) -> Cu2+(aq, 0.0010 M) + 2e-
In the half-cell with the concentrated Cu2+ ion solution, at the copper electrode, copper(II) ions gain electrons to beome copper atoms, thus decreasing the Cu2+ ion concentration in solution. The newly formed Cu atoms are platted on the electrode. This copper electrode increases in mass.
The reaction occuring at the cathode is a reduction half-reaction
Cu2+(aq, 2.00 M) + 2e- -> Cu(s)
The AACT Galvanic Cells/Voltaic Cells 2 computer simulation has animations at the particle level of these two half-reactions. In the animation below, the half-cell solution on the right should be a light blue color Cu2+ solution and the half-cell solution on the left should be a dark blue color Cu2+ solution. [A correct image will be used as soon as it becomes available.]
The cell potential, E, can be determined using the Nernst equation
The potential difference or cell potential depends upon
- the identity of the half-reactions occuring in each half-cell and
- the concentration of the ions in each of the solutions.
The potential difference (voltage) does not depend upon the amount of material of the electrode and solution (mass and volume) in the half-cells. Current and the duration of the voltage does depend upon the amount of material.
- Potential difference in an electrochemical cell does depend upon the ratio of the concentration of ions in the two half-cells.
For a copper-copper concentration cell the cell potential is related to the concentration of Cu2+ ions in solution. The greater the difference between the concentration of Cu2+ in the two half-cells, [Cu2+], the larger the potentail difference (voltage). A graph of cell potential difference versus log various [Cu2+] , dilute half-cell, shows this relationship. The concentration of Cu2+ in the other half-cell is kept at 1.00 M.
Michiel Vogelezang, and Adri Verdonk, World Journal of Chemical Education, vol. 8, no. 3 (2020).
Start the demonstration by showing the voltage of a copper-copper concentration cell with both solutions 1.00 M Cu2+. The cell potential is 0.00 Volts
Next set-up a copper-copper concentration cell one solution 2.00 M Cu2+and the other solutions 0.001 M Cu2+. The cell potential is +0.01 Volts.
The reaction occuring at the anode is an oxidation half-reaction Cu(s) -> Cu2+(aq, 0.0010 M) + 2e-
The reaction occuring at the cathode is a reduction half-reaction Cu2+(aq, 2.00 M) + 2e- -> Cu(s)
The cell reaction is Cu(s) -> Cu2+(aq, 0.001 M) + 2e-
Cu2+(aq, 2.00 M) + 2e- -> Cu(s)
___________________________
Cu2+(aq, 2.00 M) --> Cu2+(aq, 0.0010 M)
Q = Cu2+(aq, 0.0010 M)/Cu2+(aq, 2.00 M) = [Cu2+ dilute]/ [Cu2+ concentrated] = 0.0010/2.00
E = E° -(0.0592)/2 log(0.0010/2.00) = 0.00 V - (0.0592)/2(log 0.0005) = +0.097 Volts = +0.10 Volts
Solutions, Chemicals and Materials
- 2.00 and 1.00 molar copper (II) sulfate solution
- 0.10 M. 0.010 M, 10 millimolar copper(II) sulfate solution
- 20 g of solid copper(II) sulfate
- 1.00 M potassium sulfate (to maintain ionic strength in the low conentration solutions)
- 2.00 M potassium sulfate (solution for the salt bridge)
- Two clean copper metal strips
- 2.00 M and 1.00 Molar zinc (II) sulfate solution
- 20 g of solid zinc sulfate
- 0.10 M, 0.010 M, 10 millimolar zinc (II) sulfate solution
- Two clean zinc metal strips
- 200 mL deionized water
Equipment
- Four 400 mL beakers
- Two salt bridges filled with potassium sulfate
- Digital Voltmeter
- Red and black wires with alligator clips
- Label Tape
- Paper towels
Procedures
I. Copper concentration cell (same metal electrodes, same type of ions in solution, different concentration of ions) Copper electrodes, CuSO4(aq) solutions
A. Place 100 mL of 1.0 M copper (II) sulfate solution in a beaker and place 100 mL of 1.0 M solution of copper(II) sulfate in the other beaker. Place copper metal strips in each solution. Use the black and red wires with aligator clips to connect the copper strips to the voltmeter. Complete the circuit by inserting a salt bridge between the two solutions. The voltage measurement should be zero Volts.
B. Place 100 mL of 2.0 M copper (II) sulfate solution in a beaker and place 100 mL of 10 millimolar (0.001M) solution of copper(II) sulfate in the other beaker. Place copper metal strips in each solution. Use the black and red wires with aligator clips to connect the copper strips to the voltmeter. Note which electrode is connected to which terminal of the voltmerter.
Complete the circuit by inserting a salt bridge between the two solutions. The voltage measurement should be +0.10 Volts. Have students complete a diagram of this electrochemical cell, identfy ten things about an electrochemical cell. See activity sheet.
Show students the AACT electrochemical cell simulation of a copper concentration cell. Show the molecular scenes animations
C. Place 100 mL of 2.0 M copper (II) sulfate solution in a beaker and place 100 mL of 2.0 M solution of copper(II) sulfate in the other beaker. Place copper metal strips in each solution. Use the black and red wires with aligator clips to connect the copper strips to the voltmeter.
- Ask students to predict the voltage of this cell and explain their reasoning.
Complete the circuit by inserting a salt bridge between the two solutions. The voltage measurement should be ___ Volts. Ask student to explain.
- The potential difference or cell potential of an electrochemical cell depends upon
- the identity of the half-reactions occuring in each half-cell and
- the concentration of the ions in each of the solutions.
- The potential difference (voltage) does not depend upon the amount of material (mass and volume) in the half-cells. Doubling the volume of 0.100 M solution will not influence the half-cell potential difference or the voltage of the reduction potential. Doubling the size of an electrode will not influence the half-cell potential difference.
- If the concentration of ions in each half-cell solution is the same, the potentail difference will be zero Volts.
II. Zinc-copper Nernst electrochemical cell
A. Place 100 mL of 1.0 M copper (II) sulfate solution in a beaker and place 100 mL of 1.0 M solution of zinc sulfate in the other beaker. Place a copper metal strip in the copper(II) sulfate solution. Place a zinc metal strip in the zinc sulfate solution. Use the black and red wires with aligator clips to connect the copper strip and zinc strip to the voltmeter. Complete the circuit by inserting a salt bridge between the two solutions. The voltage measurement should be +1.10 Volts. This is a standard cell with E°.
B. A. Place 100 mL of 1.0 M copper (II) sulfate solution in a beaker and place 100 mL of 0.001 M solution of zinc sulfate in the other beaker. Place a copper metal strip in the copper(II) sulfate solution. Place a zinc metal strip in the zinc sulfate solution. Use the black and red wires with aligator clips to connect the copper strip and zinc strip to the voltmeter.
Ask students to predict the voltage of the cell.
Show students the AACT electrochemical cell simulation of a copper concentration cell. Show the molecular scenes animations.
Complete the circuit by inserting a salt bridge between the two solutions. The voltage measurement should be ____ Volts. This is a cell potential, E.
Have students complete a diagram of this electrochemical cell. See activity sheet.
C. Place 100 mL of 0.001 M copper (II) sulfate solution in a beaker and place 100 mL of 1.00 M solution of zinc sulfate in the other beaker. Place a copper metal strip in the copper(II) sulfate solution. Place a zinc metal strip in the zinc sulfate solution. Use the black and red wires with aligator clips to connect the copper strip and zinc strip to the voltmeter.
Ask students to predict the voltage of the cell.
Complete the circuit by inserting a salt bridge between the two solutions. The voltage measurement should be ____ Volts. This is a cell potential, E.
Have students complete a diagram of this electrochemical cell. See activity sheet.
Chemistry
In most electrochemical cells, the actual cell potential is different from the standard cell potential, E°, because the temperature of the solution and metal sree not at 1.00 atm and 25°C, and the concentrations of the ions in the solutions are not exactly 1.000 M. The (maximum) work that an electrochemical cell can do is equal to its Gibbs free energy, which for standard conditions is:
where n is the number of moles of electrons transferred in the reaction (per mole of reactant or product), F is Faraday's constant, which is the number of coulombs per mole of electrons, or 1.6022 × 10-19 C × 6.0221 × 1023/mol = 96486 C/mol, and E° is the standard cell potential. The right-hand side of this equation is the product of charge and potential, giving units of energy or work (J). For a reaction that is not in equilibrium, the Gibbs free energy is:
where Q, the reaction quotient, equals the product of the concentrations of the products, each raised to the power of its stoichiometric coefficient, divided by the product of the reactant concentrations, each raised to the power of its stoichiometric coefficient. By substituting -nFE and -nFE° into the above equation and then dividing each side by nF, the Nernst equation is obtained:
When the temperature is 298 K, using the values for R and F and multiplying by the appropriate factor to convert from base e to base 10, 2.303:
When the concentrations of both solutions are 1.00 M, under standard conditions of 25°C and a pressure of 1 atm, Q equals one, the logarithm of 1 equals zero. Under these conditions, the nonstandard cell potential, E, equals the standard cell potential, E°. When Q equals K, the equilibrium constant, that is, when the reaction proceeds until equilibrium has been reached, the second term equals the standard cell potential, and the cell potential goes to zero.
The cell potential does not depend on the how much material is in the cell (mass and volume), but does depend on the ratios of the solution concentrations raised to the appropriate stoichiometric powers. The reaction provides the energy per unit charge to an external circuit. This is a measure of how far the reaction is from equilibrium. How much material is in the cell - the concentration of ions and the mass of each electrode, are the governing factors of the total current the cell can provide.
Applying the Nernst equation to a simple electrochemical cell such as the Zn/Cu cell illustrates how the cell voltage varies as the reaction progresses and the concentrations of the dissolved ions change. The overall reaction for this voltaic cell is
Zn(s) + Cu2+(aq) -> Zn2+(aq) + Cu(s)
The reaction quotient is
Q = [Zn2+]/[Cu2+] = [2.00 M]/[0.0010M] = 2,000.
Suppose that the cell initially contains 2.0 M Cu2+ and 1.0 × 10−3 M Zn2+. The initial voltage measured when the cell is connected can then be calculated from Equation
Q = [Zn2+]/[Cu2+] = [2.00 M]/[0.0010M] = 0.0005
E = +1.10 Volts - (0.0592/2)log(0.0005)
E = +1.10 Volts - (0.0296)(-3.30) = +1.10 Volts + 0.097 = +1.19 Volts.
The initial voltage is greater than E° because Q is a small value less than 1.00 and the log of a number less than 1.00 is negative value. The left most term in the Nernst equation becomes a positive value and is added to the standard cell potential.
Starting with a large Cu2+ concentration and a small Zn2+ concentration, as the reaction proceeds, [Zn2+] in the anode compartment increases as the zinc electrode undergoes oxidation
Zn -> Zn2+ + 2e-
[Cu2+] in the cathode compartment decreases as Cu2+ ions in solution undergo reduction and copper metal is deposited on the electrode.
Cu2+ + 2e- -> Cu(s)
During the cell reaction, the ratio Q = [Zn2+]/[Cu2+] steadily increases, and the cell voltage therefore steadily decreases. The cell is not at equilibrium, shifting the equilibrium to the left goes toward equilibrium. Eventually, [Zn2+] = [Cu2+], so Q = 1 and Ecell = E°cell. Beyond this point, because zinc is the active metal a reaction will still occur. [Zn2+] will continue to increase in the anode compartment, and [Cu2+] will continue to decrease in the cathode compartment. The value of Q will increase further, leading to a continuing decrease in Ecell. When the concentrations in the two compartments are the opposite of the initial concentrations (i.e., 2.0 M Zn2+ and 1.0 × 10−4 M Cu2+), Q = 1.0 × 104, and the cell potential will be reduced to +1.00 V. A plot of Ecell versus log Q reveals a linear relationship.
Graph is from https://chem.libretexts.org/Bookshelves/General_Chemistry/Map:_Chemistry...
Zinc Concentration Cell
[image from MrGrodskiChemistry, accessed, May 2023]
In a standard electrochemical cell with two copper half-cells, Cu(s)|Cu2+(aq)||Cu2+(aq)|Cu(s) where the cell reaction is
Cu(s) + Cu2+(aq)[1.0M] → Cu2+(aq)[1.0M] + Cu(s)
Using the standard reduction potentials for the copper and zinc half cells, we find (comparing standard reduction potentials)EoCu|Cu2+||Cu2+|Cu = EoCu2+|Cu - EoCu2+|Cu = 0.34V - 0.34V = 0.0V
When the concentrations of copper ions in each half-cells are equal, the system is at equilibrium. At the point of equilibrium, there is no driving force in the reaction, the voltage is zero. EoCu2+|Cu = 0 and ΔGoCu|Cu2+||Cu2+|Cu = 0 indicates that the reaction under standard state conditions of 1.0M solutions has no driving force
K = 10nEo/0.0592V = 1 (at 25C)
In any concentration cell when the half-cells have different concentrations the driving force is to have the concentration of ions in each half-cell equal. The concentrated solution will decrease in ions and the dilute will increase in ions.
Cu(s) + Cu2+(aq)[concentrated] → Cu2+(aq)[dilute] + Cu(s)
We know that EoCu|Cu2+||Cu2+|Cu=0, n=2, and Q=[Cu2+]dilute/[Cu2+]concentrated. The cell voltage at 25C will be given by the Nernst equationECu|Cu2+||Cu2+|Cu = EoCu|Cu2+||Cu2+|Cu - (0.0592V/n) log10 Q
which becomesECu|Cu2+||Cu2+|Cu = - 0.0296V log10 [Cu2+]dilute/[Cu2+]concentrated
When the concentration of copper ion in the concentrated solution half-cell exceeds the concentration of copper ions in the dilute solution half-cell, Q < 1, the cell voltage E > 0, and ΔG < 0. The reaction will proceed spontaneously from concentrated solution to the dilute solution to increase the concentration of copper ions in the dilute solution half cell.
Some student Difficulties with Galvanic Cells or Electrochemical Cells
1. Cell potentials are obtained by adding individual reduction potentials
2. Anodes, like anions, are always negatively charged and release electrons, and cathodes, like cations, are always positively charged and attract electrons.
3. The anode is positively charged because it has lost electrons. The cathode is negatively charged because it has gained electrons.
4. Electrons flow through the salt bridge and the electrolyte solutions to complete the circuit,
Learning Objectives
References
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Arevalo, Agustin; Pastor, Gloria; Verification of the Nernst Equation. Journal of Chemical Education, 1985, Vol.62 (10), p.882.
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Greenbowe, T.J.; Gelder, J.I., Boyd, A, Wixon, M. (2020). Galvanic/Voltaic Cells 2. American Chemical Society, American Association of Chemistry Teachers, Washington, D.C. https://teachchemistry.org/classroom-resources/galvanic-voltaic-cells-2
Johnstone, A.H. (1993). "The development of chemistry teaching: A changing response to changing demand. " Journal of Chemical Education, 70(9), 701-705.
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Sanger, M. J. and Greenbowe, T.J. (1997). “Common Student Misconceptions in Electrochemistry: Galvanic, Electrolytic, and Concentration Cells.” Journal of Research in Science Teaching, 34(4), 377-398.
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Francisco J. Vidal-Iglesias, Jose Solla-Gullon,*Antonio Rodes, Enrique Herrero, and Antonio Aldaz. Understanding the Nernst Equation and Other Electrochemical Concepts: An Easy Experimental Approach for Students J.Chem.Educ. 2012, 89,936−939.
Michiel Vogelezang, and Adri Verdonk, “A Possible Electrochemical Route to a Thermodynamic Redox Reaction Equilibrium Constant in Secondary Education: An Attempt to Come from Science Fiction to Science Education?” World Journal of Chemical Education, vol. 8, no. 3 (2020): 128-140. doi: 10.12691/wjce-8-3-5.