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Simulation Activity: Galvanic/Voltaic Cells Mark as Favorite (29 Favorites)

ACTIVITY in Net Ionic Equation, Reduction, Redox Reaction, Reduction Potentials, Galvanic Cells, Oxidation, Half Reactions, Cathode, Anode, Electron Transfer, Electrons. Last updated July 25, 2023.

Summary

In this activity, students will use a simulation to create a variety of galvanic/voltaic cells with different electrodes. They will record the cell potential from the voltmeter and will use their data to determine the reduction potential of each half reaction. Students will also identify anodes and cathodes, write half reaction equations and full chemical equations, and view what is happening in each half cell and the salt bridge on a molecular scale.

Grade Level

High School

AP Chemistry Curriculum Framework

This activity supports the following units, topics, and learning objectives:

  • Unit 4: Chemical Reactions
    • Topic 4.2: Net Ionic Equations
      • TRA-1.B: Represent changes in matter with a balanced chemical or net ionic equation: a. For physical changes. b. For given information about the identity of the reactants and/or product. c. For ions in a given chemical reaction.
  • Unit 9: Applications of Thermodynamics
    • Topic 9.7: Galvanic (Voltaic) and Electrolytic Cells
      • ENE-6.A: Explain the relationship between the physical components of an electrochemical cell and the overall operational principles of the cell.

NGSS Alignment

This activity will help prepare your students to meet the performance expectations in the following standards:

  • HS-PS1-2: Construct and revise an explanation for the outcome of a simple chemical reaction based on the outermost electron states of atoms, trends in the periodic table, and knowledge of the patterns of chemical properties.
  • Scientific and Engineering Practices:
    • Developing and Using Models
    • Analyzing and Interpreting Data

Objectives

By the end of this activity, students should be able to:

  • Identify cathodes and anodes in each galvanic/voltaic cell.
  • Write half-reactions that occur in the galvanic/voltaic cells they build using the simulation.
  • Use the standard hydrogen electrode as a reference to determine the reduction potentials of each half reaction in this simulation.
  • Describe how the electrons flow through the voltaic cell to create an electric current.

Chemistry Topics

This activity supports students’ understanding of:

  • Galvanic (voltaic) cells
  • Reduction potentials
  • Redox reactions
  • Half reactions
  • Cathode and anode

Time

  • Teacher Preparation: 10 minutes
  • Lesson: 45–60 minutes

Materials

Safety

  • No specific safety precautions need to be observed for this activity.

Teacher Notes

  • Many thanks to Tom Greenbowe and John Gelder for their input on this simulation, which was inspired by their Flash-based simulation. Since Flash is no longer supported, they provided valuable insight as we designed this new simulation and resource based on their originals.
  • The Metals in Aqueous Solutions simulation simulation and the corresponding activity may be helpful to use with students prior to this one. Since the electrodes used in this simulation were also used in the Metals in Aqueous Solutions simulation, students could draw connections between the activity of various metals and their reduction potentials.
  • Most students will need to be introduced to the concepts and terms found in the “Pre-activity Questions” in class before completing this activity. Alternatively, you could have students watch a pre-recorded lecture on galvanic cells on their own and complete the definitions as a homework assignment.
  • Be sure that students are aware that the terms “galvanic cell” and “voltaic cell” are equivalent, especially if having them watch video lectures or look up definitions on their own.
  • It might be helpful for students to go through one example together as a class, especially if the concepts are relatively new to them. For example, you could build Cell #1 together (Part I), and then write the half reactions and net ionic equation for that cell (Part II). Then students could complete the rest on their own.
  • Part II has students identify half reactions and the net ionic equations for all 10 cells. If you are short on time, you could ask students to do a selection of cells, rather than all 10, or to divide and conquer (ex: half the class does the even numbered cells, and the other half does the odd numbers.)
  • Be sure students understand that a particular half reaction does not always take place at the cathode or the anode – this depends on what other electrode it is paired with and whether it has a higher or lower reduction potential.
  • In this simulation, the left electrode is always the cathode and the right electrode is always the anode. It is wise, once students get the hang of this, to have them conduct their own investigation to recognize that this isn’t always the case. (This is particularly relevant to Part I question C and Analysis questions 3 and 4.) A good activity from the AACT library is this one: Four-Way Galvanic Cell.
    • Students often have a hard time conceptualizing the salt bridge, particularly the ends being porous but not completely open. If you have a glass salt bridge in the lab, you could show them a galvanic cell set up and demonstrate how the liquid in the salt bridge doesn’t just pour out. You could also show them other types of salt bridges, such as a porous barrier or filter paper soaked in aqueous salt solution connecting the two half cells.
    • Students sometimes think electrons are moving through the salt bridge, rather than the wire. Be sure students understand that the electrons are moving along the wire from one electrode to the other, and the salt bridge provides ions that move into the solutions to keep each half cell electrically neutral as electrons move from one to the other.
    • Another common misconception students have about aqueous solutions is that they don’t realize that the ions separate in water. For example, if they see “NaCl (aq),” they picture units of “NaCl” (a sodium and chlorine atom attached) rather than separate Na+ and Cl– ions. It might be helpful before starting this activity to have students draw the particles in a beaker containing an aqueous solution so you can address this if it is a misconception held by any of your students.
      • The ions used in the salt bridge are Na+ and NO3–.
      • The standard hydrogen electrode (SHE) operates under standard conditions: 1 atm of H2 (g), 298 K, 1.0 M HCl (aq). The hydrogen gas is pumped through a length of tubing into the electrode via the inlet valve on the upper part of the electrode container. (The source of the hydrogen is not shown in this simulation for the sake of simplicity.) The piece of platinum in the electrode is a catalyst and is not consumed in the reaction, which is why that electrode, unlike the other metal electrodes, does not react. The species that are actually reacting are H2(g) and H+(aq) from the HCl solution.
      • All solutions in this simulation are standard 1.0 M solutions. The reduction potentials would differ slightly with changes in solution concentration.
      • This activity would be most useful in a unit on electrochemistry, but you could also use it when discussing redox chemistry and/or chemical reactions (specifically, single replacement reactions). You could also revisit topics from previous units and add additional layers of complexity. Other adjacent topics that are could be discussed in more detail include:
        • Activity series
        • Spontaneous and nonspontaneous reactions
        • Electrolytic cells
      • In this simulation, the molecular view only shows the metal atoms and ions, as those are the ones that have the potential to change. The anion is the same – nitrate – for all solutions in the simulation (except for the hydrochloric acid, in which case it is chloride) and doesn’t change in the single replacement reactions, so it is excluded for clarity. Similarly, water molecules are not shown as they do not change either and would far outnumber the ions in the solution. Without these spectator species shown, it is easier for students to see what changes are occurring, but you could have a discussion with students about what else is present in the beakers.
      • Related classroom resources from AACT Library that may be used to further teach this topic:
      • Students can easily access this simulation from the following link:

      For the Student

      Lesson

      Background

      Galvanic or voltaic cells harness the power of moving electrons to produce electricity as electrons move from one substance to another in a spontaneous redox reaction. These substances are the electrodes, and in this activity you will examine how using different substances as electrodes affect the overall cell potential. You will also determine the identity of the cathode and anode of each galvanic cell, identify the half reactions that occur and their reduction potentials, and describe the reaction and movement of particles at a molecular level.

      Pre-activity Questions

      Define the following terms in the space provided below:

      1. Reduction
      2. Oxidation
      3. Cathode
      4. Anode
      5. Half cell
      6. Half reaction
      7. Galvanic/voltaic cell
      8. Cell potential
      9. Reduction potential
      10. Salt bridge

      Instructions

      Go to https://teachchemistry.org/classroom-resources/voltaic-cells to access the simulation and complete the activity below.

      Part I—Building Galvanic Cells

      1. Complete Table 1 below by using the simulation to build galvanic cells with electrodes as listed in the first two columns. Record the cell potential (E°cell) from the voltmeter, the direction of electron flow from electrode to electrode (see the molecular scale views of each half cell), and the direction of ion flow from the salt bridge towards each half cell (see the molecular scale view of the salt bridge).
      Table 1. Building Galvanic cells.
      Cell # Left Half Cell Right Half Cell Cell Potential (E0cell) Direction of Electron Flow Direction of Salt Bridge Ion Flow

      1

      Ag/AgNO3(aq)

      Cu/Cu(NO3)2(aq)

      From ___ to ___

      Na+ moves towards _____, NO3– moves towards_____

      2

      Ag/AgNO3(aq)

      H2/HCl (aq)

      From ___ to ___

      Na+ moves towards _____, NO3– moves towards_____

      3

      Cu/Cu(NO3)2 (aq)

      H2/HCl (aq)

      From ___ to ___

      Na+ moves towards _____, NO3– moves towards_____

      4

      Ag/AgNO3 (aq)

      Zn/Zn(NO3)2 (aq)

      From ___ to ___

      Na+ moves towards _____, NO3– moves towards_____

      5

      Cu/Cu(NO3)2 (aq)

      Zn/Zn(NO3)2 (aq)

      From ___ to ___

      Na+ moves towards _____, NO3– moves towards_____

      6

      H2/HCl (aq)

      Zn/Zn(NO3)2 (aq)

      From ___ to ___

      Na+ moves towards _____, NO3– moves towards_____

      7

      Ag/AgNO3 (aq)

      Mg/Mg(NO3)2 (aq)

      From ___ to ___

      Na+ moves towards _____, NO3– moves towards_____

      8

      Cu/Cu(NO3)2 (aq)

      Mg/Mg(NO3)2 (aq)

      From ___ to ___

      Na+ moves towards _____, NO3– moves towards_____

      9

      H2/HCl (aq)

      Mg/Mg(NO3)2 (aq)

      From ___ to ___

      Na+ moves towards _____, NO3– moves towards_____

      10

      Zn/Zn(NO3)2 (aq)

      Mg/Mg(NO3)2 (aq)

      From ___ to ___

      Na+ moves towards _____, NO3– moves towards_____
      1. Choose one of the galvanic cells you built in Table I and reverse the left and right half cells. What changes when you do that? What stays the same? (Be sure to compare them at both the macroscopic level and the molecular view!)
      2. In this simulation, what arrangement of half cells results in a positive cell potential and what results in a negative cell potential? (Note that this holds true for this particular simulation only and is not always the case.)

      Part II—Writing Half Reactions and Net Ionic Equations

      1. For each galvanic cell you built in the simulation, write the two half reactions that occurred in each half cell in the table below. Be sure to correctly indicate which half reaction occurs at the cathode and which occurs at the anode.
      2. Then, combine them to write the full net ionic equation in the last column. (Make sure everything—including charges!—balances in the net ionic equation.)
      Table 2. Half Reactions and Net Ionic Equations.
      Cell # Cathode Half Reaction Anode Half Reaction Net Ionic Equation

      1

      2

      3

      4

      5

      6

      7

      8

      9

      10

      Part III—Determining Reduction Potentials of Half Reactions

      Cell potential (E°cell) is measured in Volts (V) and is displayed on the voltmeter in the simulation. The cell potential is the potential difference, or the difference between the reduction potentials of the cathode and anode half reactions. Since the oxidation half reaction (at the anode) is the reverse of a reduction reaction, its reduction potential is subtracted from the reduction half reaction (at the cathode) to find E°cell. This can be written mathematically as:

      E°cell = E°red,cathode − E°red,anode

      The reduction potential for each half reaction cannot be measured directly since they don’t happen in isolation, but can be determined by comparing them to the same standard. Chemists assign a value of 0.00 V to the H2/HCl half reaction.

      1. Use the galvanic cells you built in Part I that included the H2/HCl electrode (cell #2, 3, 6, 9) to determine the reduction potentials for the other half reactions. Record your answers in Table 3 below. (Use the cell potentials you recorded in Table 1 in the fourth column.) Show your work in the space below the table.

        Table 3. Reduction Potentials.
        Left Half Cell (Cathode) Right Half Cell (Anode) Cell Potential (E°cell) E°red,cathode E°red,anode

        Ag/AgNO3(aq)

        		 H2/HCl (aq)

        0.00 V

        Cu/Cu(NO3)2(aq)

        H2/HCl (aq)

        0.00 V

        H2/HCl (aq)

        Mg/Mg(NO3)2(aq)

        0.00 V

        H2/HCl (aq)

        Zn/Zn(NO3)2 (aq)

        0.00 V


      2. Choose two of the other galvanic cells you built in Part I (ones that do not involve the H2/HCl half reaction) and calculate E°cell using the values you determined above for each half reaction above. Show your work below. Do they match the voltmeter readings?
      3. More reactive elements tend to be found in compounds with other elements, whereas less reactive elements tend to be found in their pure, elemental form. Based on what you saw in the reactions that occurred in this simulation, put the elements (Zn, Cu, Ag, H2, and Mg) in order from least reactive to most reactive. Explain how you developed your list. This list is called an activity series.
      4. Look at the activity series you created in the previous question and compare it to the reduction potential values you recorded in Table 3. Can you determine a pattern that relates an element’s position in the activity series to its reduction potential relative to the other elements?

      Analysis

      1. Notice the change in shape/size of the electrodes at the macroscopic level as the reaction occurs. Which electrode (cathode or anode) ends up with more material than it started with, and which ends up with less? Explain this by drawing connections to what you saw on the molecular scale view.
      2. How would the function of the galvanic cell be affected if the salt bridge contained a very dilute solution of NaNO3 (aq)? Explain.
      3. Using the information in the diagram below and the information you obtained from the simulation, determine the reduction potential for the Pb/Pb(NO3)2 (aq) half cell. Show your work in the space to the right of the diagram.
      4. Imaginary metal X is the electrode in the left beaker and imaginary metal Z is the electrode in the right beaker, each with their corresponding nitrate solutions. The salt bridge contains NaNO3 (aq). The voltmeter reads −1.85 V. Use what you have learned from this simulation to determine the following:
        1. Which way are electrons flowing?
        2. Which is the reduction half reaction, which is oxidation?
        3. Which is cathode, which is anode?
        4. Which way are the Na+ and NO3− ions moving from the salt bridge?
        5. Which pure metal, X or Z, is more reactive?
        6. Which half reaction has the higher reduction potential?
        7. How would you determine the reduction potentials of each half reaction?