Teacher Notes

Electrochemistry Target Lab

Guided-Inquiry Wet/Dry Kit

Materials Included In Kit

Copper foil, Cu(s), 3" x 12" sheet
Copper(II) nitrate solution, Cu(NO3)2, 1.0 M, 100 mL
Iron nail, Fe(s), 12
Iron(III) nitrate solution, Fe(NO3)3, 1.0 M, 100 mL
Lead foil, Pb(s), 3" x 12" sheet
Lead nitrate solution, Pb(NO3)2, 1.0 M, 100 mL
Magnesium nitrate solution, Mg(NO3)2, 1.0 M, 100 mL
Magnesium ribbon, Mg(s), 15
Potassium nitrate solution, KNO3, 1.0 M, 200 mL
Silver foil, Ag(s), 5 g
Silver nitrate solution, AgNO3, 1.0 M, 100 mL
Zinc strip, Zn(s), 5" x ½", 10
Zinc nitrate solution, Zn(NO3)2, 1.0 M, 100 mL
Filter paper
Pipets, Beral-type, 120

Additional Materials Required

Beaker, 50-mL
Graduated cylinder, 10-mL
Reaction plate, 24-well
Sandpaper or steel wool
Scissors, Heavy-duty†
Voltmeter/multimeter
*for each lab group)
for Prelab Preparation

Prelab Preparation

  1. Cut 12, 2-cm pieces of metal foil or ribbon for each metal—copper, lead, magnesium, silver and zinc. An iron nail is used for the iron electrode.
  2. Cut the filter paper into strips approximately 4 cm long x 0.6 cm wide. Have students soak the filter paper for a few seconds in the potassium nitrate solution before use as a salt bridge. Use a fresh salt bridge for each battery.
  3. Familiarize yourself and your students with the digital voltmeter/multimeter. An analog meter may be used as well.

Safety Precautions

Silver nitrate solution is toxic by ingestion and irritating to body tissue. It also stains skin and clothing. Lead nitrate solution is a possible carcinogen. It is also moderately toxic by ingestion and inhalation, and is irritating to eyes, skin and mucous membranes. Zinc nitrate solution is slightly toxic by ingestion and corrosive to body tissue/severe tissue irritant. Copper(II) nitrate solution is slightly toxic by ingestion; it is irritating to skin, eyes and mucous membranes. Iron(III) nitrate solution is corrosive to body tissue. Magnesium nitrate solution is a body tissue irritant. Wear chemical splash goggles, chemical-resistant gloves and apron. Remind students to wash hands thoroughly with soap and water before leaving the laboratory. Please review 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 lead nitrate solution may be handled according to Flinn Suggested Disposal Method #27f. The silver nitrate solution may be handled according to Flinn Suggested Disposal Method #11. The remaining solutions may be rinsed down the drain with excess water according to Flinn Suggested Disposal Method #26b.

Lab Hints

  • Enough materials are provided in this kit for 24 students working in pairs, or for 12 groups of students. It is important to allow time between the Prelab Homework Assignment and the lab activity. Once students turn in the homework answers, figures, and procedure, check it for safety and accuracy before implementation in the lab.
  • The 24-well reaction plates are most convenient for preparing the small cells. However, small beakers may also be used.
  • When using the voltmeter/multimeter, be sure that connections are tight and that the metal electrodes are shiny. The maximum voltage will be less than 3 volts. Students may need help in setting the ideal voltage range on the meter. 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/ multimeter gives a negative voltage, reverse the connections. When a positive voltage is obtained, the electrode connected to the positive terminal is the cathode and is undergoing reduction, while oxidation is occurring at the electrode connected to the negative terminal, the anode.
  • A strip of filter paper soaked in potassium nitrate solution is used for the salt bridge. Use a fresh strip of paper for each measurement.
  • See the Example Procedure for possible voltages to assign students.

Teacher Tips

  • For the voltage challenge, you can require students to only use standard conditions (i.e., 1 M solutions). Another option, if you would like, is to allow your students to adjust the concentration of the various salt solutions. This would then lead to non-standard conditions and the voltage would change based on the changes of concentration. If students are allowed to adjust the concentrations, you should also introduce the Nernst equation. While students do not have to calculate using the Nernst equation on the AP exam, students are required to conceptually reason on how the reaction will behave based on changes to the cell.
  • Another extension for students are particulate drawings. Have students draw zoomed in views of ions turning into solid atoms with electron transfer or vice versa. Students can also draw how ions interact with water molecules.
  • Flinn Scientific has excellent video resources that enhance the teaching experience. Simply type in the keyword electrolysis, electrochemistry, or Hoffman apparatus to pull up some great videos.
  • The Colorful Electrolysis Demonstration Kit is a great extension to this lab. This demonstration kit is available from Flinn (Catalog No. AP6467).
  • Another great kit is the Electrolysis of Water Student Laboratory Kit, available from Flinn (Catalog No. AP4359).

Further Extensions

Alignment to the Curriculum Framework for AP® Chemistry 

Enduring Understandings and Essential Knowledge
Atoms are conserved in physical and chemical processes (1E)
1E1: Physical and chemical processes can be depicted symbolically; when this is done, the illustration must conserve all atoms of all types.

Chemical changes are represented by a balanced chemical equation that identifies the ratios with which reactants react and products form. (3A)
3A1: A chemical change may be represented by a molecular, ionic, or net ionic equation.

Chemical reactions can be classified by considering what the reactants are, what the products are, or how they change from one into the other. Classes of chemical reactions include synthesis, decomposition, acid−base, and oxidation−reduction reactions. (3B)
3B3: In oxidation−reduction (redox) reactions, there is a net transfer of electrons. The species that loses electrons is oxidized, and the species that gains electrons is reduced.

Chemical and physical transformations may be observed in several ways and typically involve a change in energy. (3C)
3C3: Electrochemistry shows the interconversion between chemical and electrical energy in galvanic and electrolytic cells.

Chemical equilibrium is a dynamic, reversible state in which rates of opposing processes are equal. (6A)
6A4: The magnitude of the equilibrium constant, K, can be used to determine whether the equilibrium lies toward the reactant side or product side.

Chemical equilibrium plays an important role in acid−base chemistry and in solubility. (6C)
6C3: The solubility of a substance can be understood in terms of chemical equilibrium.

Learning Objectives
1.17 The student is able to express the law of conservation of mass quantitatively and qualitatively using symbolic representations and particulate drawings.
3.2 The student can translate an observed chemical change into a balanced chemical equation and justify the choice of equation type (molecular, ionic, or net ionic) in terms of utility for the given circumstances. 3.8 The student is able to identify redox reactions and justify the identification in terms of electron transfer.
3.12 The student can make qualitative or quantitative predictions about galvanic or electrolytic reactions based on half-cell reactions and potentials and/or Faraday’s laws.
3.13 The student can analyze data regarding galvanic or electrolytic cells to identify properties of the underlying redox reactions.
6.7 The student is able, for a reversible reaction that has a large or small K, to determine which chemical species will have very large versus very small concentrations at equilibrium.

Science Practices
1.1 The student can create representations and models of natural or man-made phenomena and systems in the domain.
1.2 The student can describe representations and models of natural or man-made phenomena and systems in the domain.
1.3 The student can refine representations and models of natural or man-made phenomena and systems in the domain.
1.4 The students can use representations and models to analyze situations or solve problems qualitatively and quantitatively.
2.2 The student can apply mathematical routines to quantities that describe natural phenomena.
2.3 The student can estimate numerically quantities that describe natural phenomena.
4.2 The student can design a plan for collecting data to answer a particular scientific question.
4.3 The student can collect data to answer a particular scientific question.
5.1 The student can analyze data to identify patterns or relationships.
6.4 The student can make claims and predictions about natural phenomena based on scientific theories and models.

Correlation to Next Generation Science Standards (NGSS)

Science & Engineering Practices

Asking questions and defining problems
Developing and using models
Planning and carrying out investigations
Analyzing and interpreting data
Using mathematics and computational thinking
Constructing explanations and designing solutions
Obtaining, evaluation, and communicating information

Disciplinary Core Ideas

HS-PS1.A: Structure and Properties of Matter
HS-PS1.B: Chemical Reactions
HS-PS2.B: Types of Interactions
HS-PS3.A: Definitions of Energy
HS-PS3.B: Conservation of Energy and Energy Transfer
HS-PS3.D: Energy in Chemical Processes
HS-ETS1.B: Developing Possible Solutions
HS-ETS1.C: Optimizing the Design Solution

Crosscutting Concepts

Patterns
Cause and effect
Scale, proportion, and quantity
Systems and system models
Energy and matter
Structure and function
Stability and change

Performance Expectations

HS-PS1-1. Use the periodic table as a model to predict the relative properties of elements based on the patterns of electrons in the outermost energy level of atoms.
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.
HS-PS1-7. Use mathematical representations to support the claim that atoms, and therefore mass, are conserved during a chemical reaction.
HS-PS3-1. Create a computational model to calculate the change in the energy of one component in a system when the change in energy of the other component(s) and energy flows in and out of the system are known.
HS-PS3-3. Design, build, and refine a device that works within given constraints to convert one form of energy into another form of energy.
HS-PS3-5. Develop and use a model of two objects interacting through electric or magnetic fields to illustrate the forces between objects and the changes in energy of the objects due to the interaction.
HS-ETS1-2. Design a solution to a complex real-world problem by breaking it down into smaller, more manageable problems that can be solved through engineering.

Answers to Prelab Questions

  1. Identify the equations below as redox or nonredox.
    1. Pb(NO3)2(aq) + Na2SO4(aq) → PbSO4(s) + 2NaNO3(aq)

      nonredox

    2. N2(g) + 3H2(g) → 2NH3(l)

      redox

    3. Mg(s) + 2HCl(aq) → MgCl2(aq) + H2(g)

      redox

    4. Cu(s) + 2AgNO3(aq) → Cu(NO3)2(aq) + Ag(s)

      redox

    5. HCl(aq) + NaOH(aq) → NaCl(aq) + H2O(l)

      nonredox

  2. Of the redox equations above:
    1. Write the oxidation numbers for all the elements in the redox reactions.
    2. Identify the species that is being oxidized.
    3. Identify the species that is being reduced.
      {12386_PreLabAnswers_Equation_1}

      Hydrogen is being oxidized; it is losing electrons. Each hydrogen is losing one electron, going from a 0 oxidation state to a +1. Nitrogen is being reduced, it is changing from a 0 oxidation state to a –3. Each nitrogen atom gains three electrons.

      {12386_PreLabAnswers_Equation_2}

      Magnesium is being oxidized; it is losing electrons. Each magnesium is losing two electrons, going from a 0 oxidation state to a +2. Hydrogen is being reduced; it is changing from a +1 oxidation state to a 0. Each hydrogen is gaining one electron.

      {12386_PreLabAnswers_Equation_3}

      Copper is being oxidized; it is losing electrons. Each copper is losing two electrons, going from a 0 oxidation state to a +2. Silver is being reduced and changing from a +1 oxidation state to a 0. Each silver atom gains one electron.

  3. In lab, a student created four battery set ups with nickel metal/nickel(II) nitrate solution as half of the batteries. The other metals and solutions tested were:

    Cu(s)/1.0 M Cu(NO3)2(aq)
    Fe(s)/1.0 M Fe(NO3)2(aq)
    Al(s)/1.0 M Al(NO3)3(aq)
    Ni(s)/1.0 M Ni(NO3)2(aq).
    The student set up a Ni/Cu battery with two 10-mL beakers as shown:

    {12386_PreLabAnswers_Figure_2}
    1. Before connecting the battery to a voltmeter, the student shined small amounts of each metal using sandpaper. She also connected the two beakers with a salt bridge made of filter paper that had been soaked in a 1.0 M KNO3(aq) solution.
      1. Why is it important to use sandpaper on the metal electrode before connecting the battery?

        Metal oxides can form on the outside of the metal electrodes. Using sandpaper or steel wool allows the oxides to be removed and the pure metal to be exposed.

      2. What is the purpose of the salt bridge?

        The salt bridge allows for a complete circuit in the battery. It is a place where ions can flow while the electrons are being transferred. The salt bridge also maintains solution neutrality.

    2. Then she connected the two metal pieces to the voltmeter, using alligator clips and got a –0.62 V reading. Which metal was attached to the positive end of the voltmeter when the voltage reading was negative?

      The nickel was attached to the positive end of the voltmeter when the reading was negative.

    3. When she switched the alligator clips attached to the metals, she got a reading of +0.62 V. Label the following items in the diagram above to match the +0.62 V reading:
      1. Cathode
      2. Anode
      3. Positive connection to the voltmeter
      4. Negative connection to the voltmeter
      5. Salt bridge
        {12386_PreLabAnswers_Figure_3}
    4. When she was done collecting data, the following chart was made using the nickel electrode as the reference standard:
      {12386_PreLabAnswers_Table_2}

      Using the information above:

      1. Which ion is most easily reduced?

        Cu2+ is the most easily reduced.

      2. Which metal is most easily oxidized?

        Al(s) is the most easily oxidized.

      3. Choose the two electrodes that will give the most voltage.

        A copper/aluminum battery would give the most voltage.

        1. Which is the anode? Explain.

          Aluminum. Oxidation occurs at the anode and aluminum has the largest potential in volts for an anode.

        2. What will the voltage be? Show work.

          0.62 V + 1.38 V = 2.00 V

        3. Write a balanced net ionic equation for the reaction.

          3Cu2+(aq) + 2Al(s) → 3Cu(s) + 2Al3+(aq)

        4. Draw a galvanic cell for the reaction above using metal nitrate solutions.
          1. Include electron movement and ion movement in your drawing.
          2. Label the anode and cathode.
            {12386_PreLabAnswers_Figure_4}

Sample Data

Part 1: Zinc Reference Standard

  1. In a 24-well reaction plate, approximately 2 mL of each 1.0 M solution like below. Use a clean pipet for each solution:
    {12386_Data_Figure_4}
  2. Obtain pieces of each metal from your instructor. Polish the metals with sandpaper or steel wool so they are shiny, and insert them into the well that contains the ion of the same metal.
  3. Use fresh strips of filter paper soaked in 1.0 M potassium nitrate as salt bridges. The electrodes to be tested are:

    Ag | Ag+
    Mg | Mg2+
    Cu | Cu2+
    Pb | Pb2+
    Fe | Fe3+
    Zn | Zn2+

  4. The zinc electrode is the standard electrode. Measure the potential difference between the zinc electrode and each of the other electrodes. Record which terminal is the anode and which is the cathode for each battery. Record the data in Data Table 1.
    {12386_Data_Table_3}

    Reduction equations arranged in decreasing order of potential (using zinc as the reference standard).

    {12386_Data_Table_4}
Part 2: Voltage Challenge

Possible voltage assignments:

0.80 V
0.90 V
0.40 V
0.70 V
1.10 V

Sample cells with sample data collected from lab:
{12386_Data_Table_5}
Numbers may vary for students. Check the students own reference chart and calculations, when reviewing how the group performs on the Voltage Challenge.

References

College Board, The. 2014. “AP Chemistry Course and Exam Description, rev. ed.” NY: The College Board. Accessed June 7, 2017. http://media.collegeboard.com/digitalServices/pdf/ap/ap-chemistry-course-and-exam-description.pdf

CRC Handbook of Chemistry and Physics, 95th ed. CRC Press: Boca Raton, FL, 2014

Student Pages

Electrochemistry Target Lab

Introduction

Batteries power so many of the devices we use on a daily basis. It’s these electrochemical reactions that keep our lives and our devices charged! In this lab, you will create a half cell reduction potential data table using a zinc half-cell as your reference standard. Then your instructor will give you a voltage requirement. Using the materials available and your half-cell reduction potential data table, you’ll have to build a battery to your instructor’s specifications. See how close you can get to the actual voltage!

Concepts

  • Electrochemical cell
  • Standard reduction potential
  • Oxidation–reduction reactions
  • Half-cell reaction

Background

An electrochemical cell results when an oxidation reaction and a reduction reaction occur, and the resulting electron transfer between the two processes occurs through an external wire. The oxidation and reduction reactions are physically separated from each other and are called half-cell reactions. A half-cell can be prepared from almost any metal in contact with a solution of its ions. Since each element has its own electron configuration, each element develops a different electrical potential, and different combinations of oxidation and reduction half-cells result in different voltages for the completed electrochemical cell.

The standard reduction potential is the voltage that a half-cell, under standard conditions (1 M, 1 atm, 25 °C), develops when it is combined with the standard hydrogen electrode, which is arbitrarily assigned a potential of zero volts. A chart of reduction half-cell reactions, arranged in order of decreasing standard reduction potential, shows the relative ease of reduction of each substance listed. The more positive the reduction potential, the easier the reduction. A spontaneous cell (a battery) can be constructed if two half-cells are connected internally using a salt bridge, and externally using a metallic connector. In an electrochemical cell, the reaction listed in the standard reduction potential chart with the more positive voltage occurs as a reduction, and the reaction listed with the less positive voltage reverses and occurs as an oxidation reaction. The cell voltage can be found by adding the voltages listed in the table, with the value of the voltage for the oxidation reaction becoming the negative of its reduction reaction voltage.

As an example, consider a cell made up of copper and aluminum half-cells.

Cu2+(aq) + 2e → Cu(s) E° = 0.34 V
Al3+(aq) + 3e → Al(s) E° = –1.66 V

The copper reaction has the more positive potential and remains a reduction reaction. The aluminum reaction with the less positive (more negative) potential is reversed and becomes an oxidation reaction. Its potential is now an oxidation potential:

Al(s) → Al3+(aq) + 3e E° = +1.66 V

The reduction potential and the oxidation potential are added to find the cell voltage:

3Cu2+(aq) + 2Al(s) → 3Cu(s) + 2Al3+(aq)
E°cell = E°reduction + E°oxidation
E°cell = 0.34 V + 1.66 V = 2.00 V

A positive value for E°cell indicates the oxidation–reduction reaction, as written, is spontaneous.

A cell representation such as the following: Zn(s) | Zn2+(1.0 M) || Cu2+(1.0 M) | Cu(s) means that a cell is constructed of zinc metal dipping into a 1.0 M solution of Zn2+. The symbol “|” refers to a phase boundary. The symbol “||” indicates a salt bridge between the zinc ion solution and the copper ion solution. The second half-cell is copper metal dipping into a 1.0 M solution of copper ions. The anode is on the left (where oxidation occurs) and the cathode is on the right (where reduction occurs).

In this laboratory, a “standard” table of electrode potentials is constructed. A value of 0.00 volts is assigned to the electrode made from zinc metal in a 1.0 M solution of zinc ions. The voltage values should correlate with those found in published tables, differing only by the value of E° for the standard zinc electrode. Published standard values are measured in solutions that have very small electrical resistance. The resistance of the experimental cell will probably cause a lowering of measured values from the ideal values.

The table of standard potentials assumes that all ion concentrations are 1.0 M, gas pressures are 1 atm, and temperature is 25 °C.

Experiment Overview

The purpose of this activity is to complete the homework assignment prior to lab to promote conceptual understanding of galvanic cells. You will first review and analyze redox reactions and electrochemical cells in the homework assignment. After analyzing your homework assignment, you will then design a procedure to create your own half-cell reduction potential using a zinc half-cell as your reference standard. You will then have to construct a cell of a given voltage from your instructor. See how close your cell is to the predicted voltage!

Prelab Questions

Complete the following homework set and write a lab procedure to be approved by your instructor prior to performing the lab. Along with your procedure, you will turn in any graphs, tables or figures you were asked to create in this homework set, and answers to the questions. Use a separate sheet of paper, if needed.

  1. Identify the equations below as redox or nonredox.
    1. Pb(NO3)2(aq) + Na2SO4(aq) → PbSO4(s) + 2NaNO3(aq)
    2. N2(g) + 3H2(g) → 2NH3(l)
    3. Mg(s) + 2HCl(aq) → MgCl2(aq) + H2(g)
    4. Cu(s) + 2AgNO3(aq) → Cu(NO3)2(aq) + Ag(s)
    5. HCl(aq) + NaOH(aq) → NaCl(aq) + H2O(l)
  2. Of the previous redox equations:
    1. Write the oxidation numbers for all the elements in the redox reactions.
    2. Identify the species that is being oxidized.
    3. Identify the species that is being reduced.
  3. In lab, a student created four battery set ups with nickel metal/nickel(II) nitrate solution as half of the batteries. The other metals and solutions tested were:

    Cu(s)/1.0 M Cu(NO3)2(aq)
    Fe(s)/1.0 M Fe(NO3)2(aq)
    Al(s)/1.0 M Al(NO3)3(aq)
    Ni(s)/1.0 M Ni(NO3)2(aq).
    The student set up a Ni/Cu battery with two 10 mL beakers as shown:

    {12386_PreLab_Figure_1}
    1. Before connecting the battery to a voltmeter, the student shined small amounts of each metal using sandpaper. She also connected the two beakers with a salt bridge made of filter paper that had been soaked in a 1.0 M KNO3(aq) solution.
      1. Why is it important to use sandpaper on the metal electrodes before connecting the battery?
      2. What is the purpose of the salt bridge?
    2. Then she connected the two metal pieces to the voltmeter, using alligator clips, and got a –0.62 V reading. Which metal was attached to the positive end of the voltmeter when the voltage reading was negative?
    3. When she switched the alligator clips attached to the metals, she got a reading of +0.62 V. Label the following items in the diagram above to match the +0.62 V reading:
      1. Cathode
      2. Anode
      3. Positive connection to the voltmeter
      4. Negative connection to the voltmeter
      5. Salt bridge
    4. When she was done collecting data, the following chart was made using the nickel electrode as the reference standard:
      {12386_PreLab_Table_1}

      Using the information above:

      1. Which ion is most easily reduced?
      2. Which metal is most easily oxidized?
      3. Choose the two electrodes that will give the most voltage.
        1. Which is the anode? Explain.
        2. What will the voltage be? Show work.
        3. Write a balanced net ionic equation for the reaction.
        4. Draw a galvanic cell for the reaction above using metal nitrate solutions.
          1. Include electron movement and ion movement in your drawing.
          2. Label the anode and cathode.
  1. Design a procedure, using the zinc electrode as the reference standard instead of the typical standard hydrogen electrode as the reference. Create a chart of reduction equations arranged in decreasing order of potential.

    You will be given the following materials:

    Copper foil, Cu(s)
    Copper(II) nitrate solution, Cu(NO3)2, 1.0 M, 4.0 mL
    Iron nail, Fe(s)
    Iron(III) nitrate solution, Fe(NO3)3, 1.0 M, 4.0 mL
    Lead foil, Pb(s)
    Lead nitrate solution, Pb(NO3)2, 1.0 M, 4.0 mL
    Magnesium nitrate solution, Mg(NO3)2, 1.0 M, 4.0 mL
    Magnesium ribbon, Mg(s)
    Potassium nitrate solution, KNO3, 1.0 M, 10.0 mL
    Silver foil, Ag(s)
    Silver nitrate solution, AgNO3, 1.0 M, 4.0 mL
    Zinc strip, Zn(s)
    Zinc nitrate solution, Zn(NO3)2, 1.0 M, 4.0 mL
    Filter paper
    Sandpaper or steel wool
    Voltmeter/multimeter
    Wires and alligator clips

  2. Once you have created your zinc electrode reference standard chart, your instructor will give you a voltage requirement. You will then have to create a battery from the materials above that is as close as possible to your assigned voltage.

Safety Precautions

Silver nitrate solution is toxic by ingestion and irritating to body tissue. It also stains skin and clothing. Lead nitrate solution is a possible carcinogen. It is also moderately toxic by ingestion and inhalation; irritating to eyes, skin and mucous membranes. Zinc nitrate solution is slightly toxic by ingestion; it is corrosive to body tissue/severe tissue irritant. Copper(II) nitrate solution is slightly toxic by ingestion and irritating to skin, eyes and mucous membranes. Iron(III) nitrate solution is corrosive to body tissue. Magnesium nitrate solution is a body tissue irritant. Wear chemical splash goggles, chemical-resistant gloves and apron. Wash hands thoroughly with soap and water before leaving the laboratory.

*Advanced Placement and AP are registered trademarks of the College Board, which was not involved in the production of, and does not endorse, these products.

Next Generation Science Standards and NGSS are registered trademarks of Achieve. Neither Achieve nor the lead states and partners that developed the Next Generation Science Standards were involved in the production of this product, and do not endorse it.