Demonstrations
Demonstration 1: Voltaic Cells
Purpose
To demonstrate the ability of chemical reactions to produce electricity. The electric potential produced depends on the nature of the metal/metal ion half cell.
Use all normal safety precautions. Use accepted procedures for disposing of heavy-metal solutions.
Procedure
Cut the filter paper in the shape of a cross. Place it in the Petri dish. Add some KNO 3 solution to the center of the cross, allowing it to saturate the paper. Place one drop of each metal ion solution on each of the four corners. Attach two different metal electrodes to alligator clips wired to the voltmeter. Touch the proper electrode to its own metal ion solution. Substitute a different metal on the alligator clip until the electric potential of each of the six combinations is measured. This procedure avoids contamination and the need for a salt bridge, giving instant results.
Demonstration 2: Fruit and Vegetable Batteries
Purpose
To show that different combinations of electrodes produce different electric potentials and that some food materials can serve as electrolytes.
Materials
Safety
Do not eat these materials. Wash hands thoroughly when the demonstration is completed.
Procedure
Cut the fruit sample in half and insert two unlike metal strips into it. Connect the strips to the voltmeter with connecting wires. Read the electric potential in volts. Experiment with placing the metal strips at various distances from each other. Try other fruit/vegetables with the same metals, and other combinations of metal strips. To reinforce concepts, have students predict in each case which of the two metals will be the cathode and which will be the anode.
Demonstration 3: Electrolysis of Water in Color
Purpose
To demonstrate electrode half-reactions.
Materials
Safety
Follow normal laboratory safety precautions.
Procedure
Set up the Hoffman apparatus with its attached power source. Prepare the volume of water needed for electrolysis by adding approximately 3 g of solid sodium sulfate and a squirt of BTB. Turn on the power and observe the color changes during electrolysis. While electrolysis is occurring, show students the color of BTB in HCl (yellow) and NaOH (blue). After each arm of the apparatus has collected at least 20 mL of gas, note that the volume in one arm is about half that in the other arm. Note the color of BTB in each arm. Collect samples of gas generated at each electrode and test with glowing splints and burning splints to determine which arm contains hydrogen gas and which contains oxygen gas.
Reactions
NOTE: A U-tube set-up is preferred for seeing color differences, but withdrawing gas samples from a U-tube is difficult.
Demonstration 4: Electrolysis of Potassium Iodide
Purpose
To show electrode half-reactions involving color changes
Materials
Safety
Use normal laboratory safety precautions. The iodine produces a brown stain on skin and clothing.
Procedure
Place a Petri dish on a plastic sheet on the overhead projector. Half-fill the dish with water and dissolve about a match-head quantity of KI in it. Make carbon electrodes by using the connecting wires to connect a pencil lead to each battery electrode. Arrange the two pencil leads so that only carbon is in contact with the KI solution in the Petri dish. Add one drop of phenolphthalein. Observe the two electrodes for evidence of a reaction.
Reactions
NOTE: You can also make an electrolysis apparatus for this reaction from a 9-V battery, two pencils, two connecting wires, a 2-hole stopper, and some tape. Sharpen two pencils and push them through the stopper. Near the erasers, carve the wood away until you can make contact with the "lead" inside. (You needn't carve wood from all around the pencil to do this.) Tape the battery to the pencils. Connect each battery terminal to the pencil "lead" near the eraser via connecting wires.
This system is much less fragile than the one involving two mechanical pencil leads, although it does take more time to make. You can attach the device to a ring stand with a universal clamp on the stopper to avoid holding it during the reaction. You can show the presence of I2 by squirting a little starch solution near the anode to show the blue-black color created by the starch test. (First demonstrate the starch test to your students with a known iodine sample.) If you use the pencil lead electrolysis apparatus with a solution of SnCl2 , you can show students Sn needles that "grow" on the cathode.
Demonstration 5: Electroplating Copper
Purpose
To demonstrate industrial uses of electrolysis.
Materials
Safety
Neutralize the acidic copper sulfate solution with an appropriate bicarbonate such as baking soda or a base like soda lime before disposing in accordance with local regulations. Dilute nitric acid can be flushed down the drain with ample water. Be careful to avoid contact with either acid. Copper salt solutions are toxic.
Procedure
Clean the copper strip by dipping it into a beaker of dilute nitric acid and washing it thoroughly. Prepare a plating solution by carefully adding 15 mL concentrated sulfuric acid slowly with continuous stirring to about 200 mL saturated copper(II) sulfate solution. Attach the copper strip to the positive battery terminal using a connecting wire. Attach the object to be plated to the negative battery terminal with the other connecting wire. Place the object to be plated and the copper strip in the solution. Rotate frequently for even coating. After 3-5 min, remove the plated object and observe.
Reactions
NOTE: An interesting reversal of the reaction can be demonstrated by switching the battery leads.
Demonstration 6: Making a Simple Battery: The Gerber Cell
Purpose
To demonstrate the ability of chemical reactions to produce electricity.
Materials
Safety
Use normal laboratory safety precautions and disposal procedures.
Procedure
Clean the Cu strip by dipping it briefly in dilute HNO3 and then washing well with water. Clean the Mg strip by dipping it briefly in dilute HCl and then washing well with water. Fill the jar about 2/3 to 3/4 full of sodium sulfate solution. Wet the dialysis tubing and tie one end in a knot. Open the other end and fill to a depth slightly less than that of the jar with copper(II) sulfate solution. Place the copper strip in the dialysis tubing. Place the magnesium strip and the dialysis tubing in the jar. Then insert the stopper to hold them in place. Observe any reaction that may occur (see Figure 4).
Figure 4. Simple battery apparatus.
Attach the ends of the Cu and Mg strips to the voltmeter with the connecting wires.
Reactions
NOTE: If you make six of these cells and hook them in series (anode to cathode), you can produce enough electric potential to operate a small radio that normally uses a 9-V battery.
Reference Summerlin and Ealy. (1985). Chemical demonstrations, Vol. 1. Washington, DC: American Chemical Society.
1. What is the role of electrons in oxidation-reduction reactions? [Electrons are the exchange particles of oxidation-reduction, moving from the oxidized to reduced species.]
2. How can chemical reactions be used to produce electricity? [If the half cells of a redox reaction are physically separated but connected by a salt bridge (to complete the internal circuit), electrons will flow through an external circuit of wire. The anion and electron movements are in opposite directions.]
3. How can electricity be used to drive a chemical reaction? [If a source of electric potential (such as a battery) is attached to electrodes in a solution or ionic melt, the applied potential forces electrons through the system. The electrode connected to the negative terminal holds an excess of electrons, attracts cations, and reduces them. This electrode is the cathode. Simultaneously, oxidation is driven at the anode as anions are attracted, lose their excess electrons to the electrode, and complete the electrical circuit. Electrolysis and electroplating processes rely on such reactions.]
4. How do materials compare in their relative ability to accept or donate electrons? [Materials tend to be classified as electron donors if they are metallic and receivers if they are nonmetallic. However, any particular reaction depends on the relative tendencies of the materials. Thus we rank materials according to their reduction potential. Even though hydrogen in contact with hydrogen ions is assigned a half-cell potential of zero volts, it (like other half cells) can either involve oxidation or reduction. The process observed depends on the nature of the other reactant. If it has a higher reduction potential, hydrogen molecules will be oxidized; if not, hydrogen ions will be reduced.]
5. How are electrochemical processes used in business and industry? [As sources of energy (electrochemical cells) and to drive chemical processes (electrolytic cells).]
6. How are electrochemical devices used in everyday life? [Applications of electrochemical cells and batteries (groups of cells in series) are myriad. Some examples include flashlights, portable appliances, pacemakers, automobile cranking systems, and submarine propulsion units.]
Counterintuitive Examples/Discrepant Events
1. Iron Can Be Protected from Corroding by Attaching It to Zinc
Iron is frequently used in construction and will corrode (oxidize) unless it is protected. One way to protect iron is to paint it. However, in many applications paint alone will not work because it is easily chipped or scratched. An effective way to protect iron is to attach a piece of zinc to it. This is frequently done to protect buried fuel tanks and hulls of ships. Zinc is easier to oxidize than iron and will lose its electrons (corrode) instead of iron. In such an application, zinc is a sacrificial cathode.
Have students observe the corrosion of 6-penny iron nails. Start with two identical 6-penny nails. Clean the surfaces of both nails with a piece of sandpaper. Wrap one nail with 10 cm of zinc wire, making sure to establish good physical contact between the two metals. Immerse nails in a beaker containing a 10% sodium chloride solution. Observe the nails over several days. The bare nail will show signs of corrosion while the one wrapped in zinc will be protected from corrosion. (Students should recognize that galvanized construction materials are made of iron protected by zinc.) This activity is useful for initiating discussion about electrochemistry prior to Activity 1 in this module.
1. Imagine an oxidation-reduction reaction as involving a hungry atom (draw a circle with chomping teeth in open mouth) eating an electron being fed to it by a nonhungry atom (closed mouth.) A voltaic cell just does the same thing, but the "hungry atom" has to be fed through a tube.
2. Use water analogies for electron transport.
Figure 5. Electron transport/ water analogies.
3. Imagine that the quantity of work that can be performed by a galvanic cell is similar to the work obtainable from a water wheel powered by falling water.
The height of the waterfall is analogous to the cell potential. The higher the waterfall, the more potential energy it has. In the case of a voltaic cell, the higher the cell potential, the more the "driving force" for an electron. The quantity of water flowing over the waterfall is analogous to the current generated by the voltaic cell. The quantity of work that can be performed by a water wheel (equivalent to the voltaic cell) depends on both how much water flows each second (equivalent to current) and the height that water falls (equivalent to the electric potential.)
Figure 6. Water wheel analogy.height
Memory Aids: Mnemonics, etc.
1. Use a cat picture with + signs for eyes to remember Cation is positive. Or use cat paw with + sign in the pad print to remind students that cations are "pawsitive."
2. LEO the lion says GER. LEO = Loss of Electrons is Oxidation. GER = Gain of Electrons is Reduction.
3. Cathode and Reduction both start with consonants; Anode and Oxidation both start with vowels.
4. "An Ox and Red Cat" (ANode involves OXidation and REDuction involves CAThode)
5. OIL RIG = Oxidation Is Loss (of electrons), Reduction Is Gain (of electrons)
 |
TABLE OF CONTENTS | TOPIC OVERVIEW | CONCEPT/SKILLS DEVELOPMENT | LINKS/CONNECTIONS | EXTENSIONS |  |
|---|