Activity 1: A Study of Voltaic Cells
Among several major types of chemical reactions, one of the most important is oxidation-reduction. You may have studied the electron transport system and photosynthesis in a biology course, for example. Both systems are just a series of oxidation-reduction reactions, each involving a species that loses electrons (is oxidized), while another species gains electrons (is reduced).
Even though oxidation cannot occur without reduction and vice versa, it is often useful to consider oxidation-reduction reactions in two parts called half-reactions. Added together, the two half-reactions make up the overall oxidation-reduction reaction.
To decide whether a particular oxidation-reduction reaction will occur, it is helpful to think about the reduction potential. Reduction potential can be considered the driving force for a half-reaction to undergo reduction--that is, to go in the direction of the species gaining an electron. Consider a spontaneous oxidation-reduction reaction. The half-reaction with the larger reduction potential will occur as a reduction reaction. The half-reaction with the smaller reduction potential will run in the reverse direction--as an oxidation reaction in which electrons are lost rather than gained. For a more complete discussion of oxidation-reduction and half-reactions, refer to your textbook.
The standard way to compare half-reactions is in a voltaic cell. A voltaic cell is an electrochemical device that can produce electrical energy from spontaneous oxidation-reduction reactions. All electrochemical cells have two electrodes--a cathode and an anode. (An electrode supplies or accepts electrons from a chemical reaction.) Reduction reactions always occur at the cathode and oxidation reactions always occur at the anode. One easy way to avoid confusion is to remember that "oxidation" and "anode" both start with vowels, whereas "reduction" and "cathode" both start with consonants. In voltaic cells, the cathode is charged positively and the anode is charged negatively. The identity of the cathode and anode is determined by the relative reduction potentials of the half-reactions that make up the voltaic cell. The electrode in the half-reaction with the more positive reduction potential is always the cathode in a voltaic cell. The electrode in the half-reaction with the less positive reduction potential is always the anode in a voltaic cell.
Unlike reactions that occur in test-tubes when substances are mixed together, voltaic cells are arranged so that half-reactions are physically separated, often into different containers called half-cells. The half-cells are connected by a conducting wire between the two electrodes. The conducting wire is often called the external circuit. An electrical current passes through the external circuit. Electrical quantities such as current (rate of flow of electrons), electric potential (electric potential energy difference between the two half cells), and resistance can be measured in this external circuit. The unit for measuring electric potential is the volt, thus electric potential is often referred to as voltage. The electricity flowing through the external circuit can be used to provide energy to systems such as flashlight lamps, audio recorders, and science experiments. A system in the external circuit is often called a load.
Figure 2. Voltaic cell.
The half-cells are also connected by an internal circuit, often provided by a salt-bridge. The salt-bridge is composed of positive and negative ions that are free to move from one half-cell to the other but do not participate in the oxidation-reduction reaction. The salt-bridge is needed to keep charge from building up in the two half-cells. When electrons flow from a half-cell through the external circuit, negative ions travel to that half-cell to maintain electrical neutrality. Conversely, when electrons flow into the other half-cell, positive ions travel to that half-cell to ensure electrical neutrality there also.
Major Chemical Concepts
1. Electrochemistry represents a subset of oxidation-reduction.
2. Electrochemical cells include electrolytes, electrolyte bridge, electrodes labeled as anode and cathode, and an external circuit through which electrons flow.
3. Chemical energy is converted into electrical energy in a voltaic cell.
4. Half-cell potentials are relative, defined by arbitrary standards and an arbitrary zero value. (To successfully complete the activity and understand the concepts, it is not necessary to introduce the hydrogen half-cell as the arbitrary zero, but you may wish to do so.)
5. Charge is conserved in cells by the movement of ions.
6. Applications of the concepts developed in the activity include batteries and protection of artifacts from corrosion.
7. Cell potential or cell voltage--voltage associated with an electrochemical cell. It can be calculated from standard potentials and the Nernst equation.
8. The volt is the unit of electrical potential. It is defined as a Joule per Coulomb (1 V = 1 J/C).
If you decide to have students calculate the half-reaction reduction potentials, this activity is most appropriate for general chemistry students. If students only measure the voltage, the activity may also be used with basic level students.
Expected Student Background
The following concepts are prerequisites for successfully understanding concepts developed in this activity: oxidation and reduction; the general nature of energy including potential energy, chemical energy, and electrical energy; the nature and behavior of charged particles, conductivity in ionic solutions; spontaneity of chemical reactions; arbitrary standards for measurement. Students should have developed fundamental laboratory skills. If you elect to have students calculate reduction potentials of the half-reactions tested, students should have arithmetic and/or calculator skills.
If you prepare salt bridges, and dropper bottles of solutions in advance (see Advance Preparation), students should complete collecting data and clean up in 30-40 min. Very efficient students will be able to complete collecting data in considerably less time. Pre- and post-laboratory discussion times will vary, depending on how much concept development is tied to the laboratory experience.
See safety instructions in the student instructions. Similar laboratory activities in commercially available curricula often use lead or nickel half-reactions. Because lead compounds are toxic and some nickel compounds have been identified as carcinogens, we recommend that these not be used.
Materials (For 24 students working in pairs)
In addition to 0.1 M solutions for students, prepare 0.5 M and 1.0 M copper(II) and zinc sulfate solutions for post-laboratory discussion. Also have available some test-tubes and strips of Zn and Cu. Use the metals and solutions recommended. Not all combinations work well under the conditions of this laboratory procedure.
Check to ensure that all materials are available. Most materials are available in high school chemistry stockrooms. Silver wire can be ordered from a chemical supply company or may be available locally at a jeweler supply house.
Prepare in advance the salt bridges. Materials needed include potassium nitrate, burner, razor blade, and thin-stemmed polyethylene pipets.
Quickly wave the pipet through burner flame to warm it. (You will be surprised at how quickly the tube softens.) While it is still soft, bend the tube into a "U" shape and allow it to cool. Do not kink the tube while bending it. Use a razor blade to cut off the bulb and excess stem. Leave about 2 cm of tube on each "arm" of the bridge. Make all the "U" tubes you will need before proceeding to the next step.
Mix 2 g potassium nitrate in 10 mL of water. Add 0.1 g agar. Boil the solution for 3 to 5 min. While the agar is cooking, stretch the end of an uncut, unbent polyethylene pipet to reduce the diameter of its tube. Do this by warming in the burner flame and stretching. Cut the end of the pipet so it is even. Remove the agar from the heat. Use the modified pipet to fill each "U" tube with the warm agar solution. Be certain to eliminate all air pockets. After the salt bridge has cooled, trim the ends with the razor blade. Store the salt bridges in a jar containing 2 M potassium nitrate. The salt bridges will last indefinitely if they do not dry out. Make several extras; those used with silver and tin half-reactions will plug up after several uses due to precipitation of silver chloride. This can be avoided by substituting tin(II) nitrate or acetate for tin(II) chloride, but the former are not commonly available in high schools.
If you do not wish to make permanent salt bridges, filter paper strips or pieces of cotton twine soaked in potassium nitrate solution work fine as temporary and easily disposable salt bridges.
Pre-laboratory work will take two or three class periods if major concept development is tied to the activity. If your students have studied in advance all concepts to be developed by the activity, use the activity for concept review. The pre-laboratory would thus focus on demonstrating use of the apparatus. We recommend, however, that the activity be used for concept development.
Review prerequisite concepts briefly to ensure student understanding (see Expected Student Background). Ask certain students to explain each concept, followed by asking other students to elaborate, agree or disagree, correct, and provide examples. Your role is to serve as a "traffic director" for questions and responses, and to record on the board or overhead appropriate ideas to structure the review. You might administer an ungraded pretest quiz a day or so before the pre-laboratory discussion to help focus discussion. One useful way to do this is to provide a list of concept words in the domain being tested, such as oxidation-reduction or energy--and ask students to draw a concept map. (Concept maps are also useful after this discussion to summarize the review.)
Once prerequisite concepts have been reviewed, introduce laboratory equipment to be used. Demonstrate how voltmeter works by testing a small battery. Show students a 24- well plate with two solutions, metal wires or strips, and a salt bridge; connect the voltmeter.
If time permits, develop, via student discussion, the most efficient arrangement of wells for testing the various combinations and the minimum number of tests required for answering the questions. As students provide suggestions, ask them to explain how their arrangement will work and how data from their suggested tests can help answer the questions. Encourage students to propose tests that do not match up the silver and tin half-reactions. Although we recommend that students develop the well arrangement and test sequence, the arrangement in Figure 3 has been found to be very efficient. You can give it to students in advance to save time, but development of thinking skills will be compromised.
As students conduct tests, have them identify the cathode and the anode, telling you what observations support their answers. Ask them whether oxidation or reduction is occurring in a particular half-cell, to define whether this is gain or loss of electrons, and to write the predicted half-reaction equation. You may wish to ask only one question of each group as you move from group to group. In that event, you might wish to prepare a checksheet matrix of question type by group so you can ask each group a conceptually different question each time you stop by.
Anticipated Student Results
Once laboratory errors such as placing the wrong metal into a well solution are taken into account, the half-reaction reduction potential order should be the same as that predicted by a standard table of reduction potentials. The actual values in volts for half reactions may differ somewhat from table values, depending on voltmeter impedance, meter reading errors, cleanliness of metal strip and wire surfaces, quality and size of salt bridge, and local concentrations of ions near the electrodes. Deviations from standard conditions under which tabled potentials are expressed should not much affect the results. (To understand why, review applications of the Nernst equation to cell potentials, as developed in college general chemistry textbooks.)
Answers to Questions
Data Analysis and Concept Development
1. See Anticipated Student Results.
2. See Anticipated Student Results.
3. Reasonable hypotheses include incorrect meter readings; using wrong metals in well plate solutions; and deviation from standard conditions of temperature, solution concentration, and/or pressure. Less likely to be proposed by students are surface coating of electrodes leading to changes in reduction potential, and lack of stirring in wells leading essentially to zero concentration of ions in well solutions near surfaces of electrodes. Accept any reasonable hypotheses as long as they are, in principle, testable.
4. For charge to be conserved in a half-cell, negative ions must replace lost electrons and positive ions must offset gained electrons.
5. Yes; No. (Be prepared to demonstrate this during post-laboratory discussion.)
Implications and Applications
1. Of metals tested in this activity, Zn and Mg can be used to protect Fe. For economic reasons, Mg is not used.
2. The Zn undergoes oxidation more easily than does Fe. (It has a larger negative--therefore lower--reduction potential.) Therefore, when Zn is attached to the Fe hull, Zn will corrode via oxidation rather than the Fe.
Ask each group to report its results. Even if four of five groups obtain the same result, it is still possible that the four made a common mistake and the one completed the test correctly. Ask students to propose hypotheses to explain any discrepancies. Ideally, student hypotheses can be tested directly at that time. A common hypothesis is that concentration makes a difference, so you should have available several solutions at higher concentrations to test that notion. By the close of this discussion, there should be general agreement on the observationsÑthe cathode and anode in various combinations and the electric potential produced by those combinationsÑ and reasons for any discrepancies. This discussion is an excellent opportunity to develop scientific thinking skills among your students.
Once observations are agreed on, process the information as a class. Repeat the questioning until an agreed order of reduction potential is established. If the activityÕs procedure is followed carefully, this order should agree with the sequence found in a table of standard reduction potentials. As part of this discussion ensure by questioning and explaining that students understand the concept of reduction potential.
After establishing the reduction potential order, give students the value from a standard table for one half-reaction used in the activityÑany half-reaction potential will do. Once students know the half-cell potential for one half-cell, they can calculate the half-cell potential for another half-cell by finding the potential measured in the laboratory for that particular half-cell combination. Because a cellÕs electric potential is the sum of the potentials of its two half-cells, students can calculate a half-cell potential for the unknown half-cell. They can then proceed to calculate half-cell potentials for each half-cell studied in the activity. Our experience indicates that students learn to solve cell-potential problems much more readily by applying the relevant concepts to the laboratory first, followed by practicing textbook-type problems, rather than the reverse.
Have students use their text and school library resources to find out what materials are used in commercially available batteries. Then ask students to use the table of half-reaction reduction potentials to explain why many batteries provide 1.5 V of electric potential. Radio Shack sells a book explaining the makeup of many common types of batteries. In addition to exploring radio batteries, students can also calculate the electric potential of car batteries based on half-cell potentials and the number of cells in a car battery. Additional activities in commercially available laboratory books are excellent for reinforcing the concepts developed in this activity. ÒExperiment 46: CorrosionÓ (Wilbraham et al., 1987) is highly recommended, as is the essentially identical activity from Carmichael and Haines (1987).
Assessing Laboratory Learning
1. Provide students with two unknown metals and matching solutions. Ask them to determine which has the higher half-reaction reduction potential and/or the value of the cell electric potential. (The former task is more difficult conceptually.) The advantage of this assessment is its direct relation to laboratory procedures used by students. Disadvantages include set-up time, time needed for students to perform, and the large quantity of equipment required if individual students are to be tested simultaneously.
2. Provide students with a group of metals, test-tubes, and solutions. Without creating voltaic cells, have them determine the order, from lowest to highest, of half-reaction reduction potentials. Advantages include ease of set up, quick implementation by students, and the conceptual nature of the task. The primary disadvantage is that it does not duplicate the studentsÕ original laboratory procedure.
If you wish, you can use questions recommended for teacher/student interactions during the activity as an evaluation tool. If you wish to assess individual student progress, make a student-by-question matrix to ensure that you ask each student the same questionsÑor at least the same number of questions. Then, as you circulate through the laboratory, ask each student a different question, checking their relative success. By circulating back through the laboratory several times, you can ask each student several questions before the end of the laboratory period, referring to your matrix to see which questions a particular student has been asked. If student performance on questions is to be used for grading, inform students in advance. We recommend that questions asked during the activity not be used for grading purposes.
Pencil and Paper
1. Provide students with written observations related to a similar laboratory activity. Ask students to interpret the information. Advantage: Duplicates the laboratory activity. Disadvantage: Requires considerable reading and therefore takes even good students a long time to complete.
2. Solve cell potential problems for half-cell potential, given the cell potential and the half-cell potential for one half-reaction. Advantage: Duplicates the laboratory activity. Disadvantage: Should probably not be used with basic level students.
3. Use questions recommended for teacher-student interactions during the laboratory activity later as written questions.
In the assessments described above, you may elect to allow students to refer to their notes, laboratory reports, and/or textbook.
Other Laboratory Activity Ideas
Give students some bubble solution made of about three parts good quality liquid detergent such as Joy and about one part glycerol, a 9-V battery, some connecting wires, and several pieces of aluminum foil. Have students electrolyze the soap solution with the Al electrodes, holding the electrodes close together so that bubbles form containing a mixture of hydrogen and oxygen gas. Holding a lighted match to the bubbles gives a satisfying but safe explosion. [Edge, 1984]
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