Materials Included In Kit

Additional Materials Required

Safety Precautions

The copper(II) sulfate solution is harmful if swallowed and causes serious skin and eye irritation. The sodium sulfate solution may be harmful if in contact with skin. Magnesium ribbon is a flammable solid. Wear chemical splash goggles, chemical-resistant gloves and a chemical-resistant apron. Wash hands thoroughly with soap and water before leaving the laboratory. Please follow all laboratory safety guidelines. Remind students to wash their hands thoroughly with soap and water before leaving the laboratory. Please review current 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. You may save all materials, including solutions, for future labs. Any leftover copper(II) sulfate and sodium sulfate solutions may be flushed down the drain with excess water according to Flinn disposal method #26b. Treated filter papers may be placed in the trash according to Flinn disposal method #26a.

Lab Hints

• An alternative to pre-cutting the filter paper and submerging into the sulfate solutions is submerging the entire filter paper in the sulfate solutions then wait until it is dry and cut to appropriate size.
• Copper tape was applied to the positive LED terminal of the battery as the conductive material to prevent reaction between Mg(s) and the copper(II) sulfate solutions with the bare LED terminal.

Teacher Tips

• Flinn Scientific has excellent video resources that enhance the teaching experience! Simply type in the key word electrolysis or Hoffman apparatus to pull up some great videos.
• The Colorful Electrolysis Demonstration is a great extension to this lab! This demonstration kit is available from Flinn (Catalog No. AP6467).

Answers to Prelab Questions

1. Define the terms oxidation and reduction.

Oxidation is a term used to describe when a substance loses electrons. Reduction describes a process in which a substance gains electrons.

2. Identify the following oxidation reaction and reduction reactions.

Cu²⁺(aq) + 2e^– → Cu(s)   Reduction
Cu(s) → Cu²⁺(aq) + 2e^–   Oxidation

3. Identify the oxidizing and reducing agent in the following reaction.

2Al(s) + 3CuCl₂(aq) → 2AlCl₃(aq) + 3Cu(s)
Aluminum metal is the reducing species and Cu²⁺ is the oxidizing species.

4. Describe the similarities and differences between galvanic cells and electrolytic cells.

Both types of cells function by oxidation–reduction reactions; a species is oxidized while another reduced. However, an electrolytic cell requires external power for a non-spontaneous reaction to occur whereas a galvanic cell’s reaction occurs spontaneously.

Answers to Questions

1. What type of cell did you build? Galvanic or electrolytic? Explain.

A galvanic cell was built because a spontaneous chemical reaction occurred (LED lit) without the need of an external power source.

2. Draw a diagram of your battery. Label all of the parts.

3. What reaction is occurring at the cathode? At the anode? Refer to a standard reduction potentials table in your textbook.

Cu2+(aq) + 2e– → Cu(s) (reduction, cathode)
Mg(s) → Mg2+ + 2e– (oxidation, anode)

4. Explain any observations noted from Part II, step 5.

Dark spots were seen on the copper(II) sulfate filter paper, which was copper metal. Gray discoloration from the magnesium metal was seen on the filter papers. This is due to the magnesium metal reaction with DI water. Without the conductive tape, the LED terminals can darken due to oxidation of Mg(s).

5. How does your battery compare to a lithium ion battery?

Lithium ion batteries are more complex, however, there are similarities to the simple battery built in lab. Both have a cathode, anode and salt bridge (electrolyte). In a lithium ion battery, the cathode is lithium and the anode is carbon. Ions first flow from anode to cathode and then from cathode to anode; this feature makes the lithium ion battery rechargeable.

References

Eggen, P.; Skaugrud, B. An Easy-to-Assemble Three-Part Galvanic Cell. J. Chem. Educ. 2015, 92(6), 1053–1055.

Introduction

Assemble your very own hand-held, tiny battery with this student lab kit! Batteries have tremendous impacts in our everyday lives. Delve into galvanic and electrolytic cells—how are these types of cells related to batteries? Start making connections. A few simple materials are provided in order to successfully complete the lab and the procedure guides you to ensure success. Enjoy this fun learning experience and take pride in your battery creation!

Watch the introduction video.

Concepts

• Half-cell reaction
• Oxidation–reduction reaction
• Galvanic vs. electrolytic cell
• Standard reduction potential

Background

In a galvanic cell, a spontaneous chemical reaction releases energy in the form of electricity (moving electrons). The chemical reaction that generates electricity in a battery is known as an oxidation–reduction reaction. Oxidation is a term used to describe when a substance loses electrons. Reduction describes a process in which a substance gains electrons. When a substance is oxidized and loses electrons, the resulting oxidized species becomes more positive. In a typical battery, the oxidized substance is converted from a neutral metal atom into a metal cation, or an ion with a positive charge. During reduction, a substance gains electrons and becomes more negative. The reduced substance is a metal cation that gains electrons to become a neutral metal atom. As an example, consider a cell made up of copper and aluminum half-cells. The copper reaction is an example of a reduction reaction and the aluminum reaction is the oxidation reaction.

Cu²⁺(aq) + 2 e^– → Cu(s)   Reduction
Al(s) → Al³⁺(aq) + 3 e^–    Oxidation

See Figure 1 for an example of a typical galvanic cell. Figure 1 can be formally represented by the following example cell notation: Zn(s) | Zn²⁺(1.0 M) | | Cu²⁺(0.0010 M) | Cu(s). In this notation 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 0.0010 M solution of Cu²⁺. The anode is on the left (where oxidation occurs) and the cathode is on the right (where reduction occurs).

While electrolytic cells function in a similar manner to galvanic cells, in contrast, an electrolytic cell requires external energy for non-spontaneous oxidation–reduction reactions to occur. Figure 2 is an example of an electrolytic cell where an electric current is passed through an aqueous solution containing an electrolyte sodium sulfate (Na₂SO₄).

The water molecules break apart or decompose into their constituent elements, hydrogen and oxygen. The overall reaction occurs as two separate, independent half-reactions. Reduction of the hydrogen atoms to elemental hydrogen (H₂) occurs at the cathode (–), while oxidation of the oxygen atoms in water to elemental oxygen (O₂) occurs at the anode (+). Each half-reaction is accompanied by the production of OH^– or H⁺ ions as shown:

Cathode: 4e^– + 4H₂O → 2H₂(g) + 4OH^–
Anode: 2H₂O → O₂(g) + 4H⁺ + 4e^–

Both types of cells are incredibly useful in the real world around us. We can generate electrons to power our devices via spontaneous reactions generated by galvanic cells. Or decompose compounds with non-spontaneous reactions accomplished by electrolytic cells. This lab activity will challenge you to build your own tiny handheld battery. Can you identify what type of cell you built? Galvanic or electrolytic?

Experiment Overview

The purpose of this activity is to build a working battery from a few simple materials. Follow the instructions to prep the materials for your battery, then follow the guided procedure for assembly. An illuminated LED indicates successful completion of the activity!

Materials

Prelab Questions

1. Define the terms oxidation and reduction.
2. Identify the following oxidation reaction and reduction reactions.

Cu²⁺(aq) + 2e^– → Cu(s)
Cu(s) → Cu²⁺(aq) + 2e^–

3. Identify the oxidizing and reducing agent in the following reaction.

2Al(s) + 3CuCl₂(aq) → 2AlCl₃(aq) + 3Cu(s)

4. Describe the similarities and differences between galvanic cells and electrolytic cells.

Safety Precautions

The copper(II) sulfate solution is harmful if swallowed and causes serious skin and eye irritation. The sodium sulfate solution may be harmful if in contact with skin. Magnesium ribbon is a flammable solid. Wear chemical splash goggles, chemical-resistant gloves and a chemical-resistant apron. Wash hands thoroughly with soap and water before leaving the laboratory. Please follow all laboratory safety guidelines.

Procedure

Part I. Battery Materials Preparation

1. Gently polish both negative and positive terminals on the LEDs with the sand paper (see Figure 3).
2. Make sure the copper conductive adhesive tape pieces and magnesium metal pieces are 2 cm in length. Cut the pieces with scissors if necessary.
3. Cover the positive terminal of each LED (see Figure 3) with the 2 cm piece of the adhesive conductive tape. Sticky side should come in contact with positive terminal.
4. Cut square shaped filter paper in the following sizes:

1. Cut 2 pieces to be submerged in copper(II) sulfate solution in about ½ cm² in size.
2. Cut 2 pieces to be submerged in the sodium sulfate solution in about 1 cm² in size.

5. Submerge filter papers from 4a into a weigh boat with 5 mL of copper(II) sulfate solution and filter papers from 4b into 5 mL of sodium sulfate solution.
6. Once the filter papers are completely coated with the designated solution, remove the papers with tweezers, place on a paper towel, and allow to completely dry. Meanwhile, read through Part II of the activity.

Part II. Battery Assembly Challenge

1. Place all of the materials on the lab bench top. Materials include prepped LEDs, magnesium metal pieces and 2 each of copper(II) sulfate and sodium sulfate filter paper squares.
2. Define the following battery components as belonging to the cathode, anode or salt bridge. How is identifying this information useful? Hint: It may be helpful to look at a standard reduction potentials table in your textbook. Which half-reaction has the more positive potential? Which the most negative?

1. Magnesium metal:
2. Copper(II) sulfate filter paper:
3. Sodium sulfate filter paper:

3. Add a drop of DI water to each filter paper. Be cautious, do not add more than 1 drop of water.
4. Contact between all battery components must be made in order to successfully light the LED. This can be accomplished by squeezing all components between the positive and negative terminals with your gloved thumb and index finger. Attempt assembling the battery now.
5. After you have correctly arranged the battery components to light the LED, disassemble the battery components and make observations. Have any of the components changed in physical appearance?
6. If the LED did not light, think about what you might do differently. In particular, think of how the components (i.e., the magnesium metal and pieces of filter paper) must be arranged (or stacked) in order for current to move through the cell. Is everything exactly where it should be?
7. (Optional) Watch the conclusion video.

Student Worksheet PDF

14164_Student1.pdf