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Making and Testing a Simple Galvanic Cell.

Making and Testing a Simple Galvanic Cell.

Introduction: (Initial Observation)

Galvanic cell is a device in which chemical energy is converted to electrical energy. In a galvanic cell Oxidation-Reduction chemical reactions produce electricity that can be used to turn on a flashlight lamp. Most dry batteries are actually some kind of galvanic cell.

A simple electrochemical cell known as Daniell cell can be made from copper and zinc metals with solutions of their sulfates. English chemist John Frederick Daniell developed this voltaic cell in 1836.

In this project I will make a simple galvanic cell and use it to experiment and identify the conditions that affect the production of electricity in a voltaic cell such as Daniell cell.


This project guide contains information that you need in order to start your project. If you have any questions or need more support about this project, click on the “Ask Question” button on the top of this page to send me a message.

If you are new in doing science project, click on “How to Start” in the main page. There you will find helpful links that describe different types of science projects, scientific method, variables, hypothesis, graph, abstract and all other general basics that you need to know.

Project advisor

Avoid contact with chemicals.
Goggles must be worn throughout experiment.

Information Gathering:

Gather information about different galvanic cells or electrochemical reactions that produce electricity. Read books, magazines or ask professionals who might know in order to learn about the factors that affect the production of electricity in a galvanic cell. Keep track of where you got your information from.

Make yourself familiar with the terms anode, cathode, oxidizing agent, reducing agent, and electromotive series.

Following are samples of information that you may find.

A Galvanic cell is an electrochemical cell that uses a spontaneous chemical reaction to produce electrical energy. Also known as a voltaic cell.

A galvanic cell is usually made of an electrolyte and two electrodes made of two different metals (with two different reactivity rates).

The primary difference between metals is the ease with which they undergo chemical reactions. The elements toward the bottom left corner of the periodic table are the metals that are the most active in the sense of being the most reactive. Lithium, sodium, and potassium all react with water, for example. The rate of this reaction increases as we go down this column, however, because these elements become more active as they become more metallic.

Classifying Metals Based on Activity

The metals are often divided into four classes on the basis of their activity, as shown in the table below.

Common Metals Divided into Classes on the Basis of Their Activity

Class I Metals: The Active Metals
Li, Na, K, Rb, Cs (Group IA)
Ca, Sr, Ba (Group IIA)
Class II Metals: The Less Active Metals
Mg, Al, Zn, Mn
Class III Metals: The Structural Metals
Cr, Fe, Sn, Pb, Cu
Class IV Metals: The Coinage Metals
Ag, Au, Pt, Hg

The most active metals are so reactive that they readily combine with the O2 and H2O vapor in the atmosphere and are therefore stored under an inert liquid, such as mineral oil. These metals are found exclusively in Groups IA and IIA of the periodic table.

The most common voltaic cells are made of two electrodes, one from class II metals and the other from class III metals. Electrolyte is usually a salt solution of the class III metal. In this way the class III metal will be reduced to metal and the class II metal will be oxidized and become a salt and enter the electrolyte solution.

For example if you use magnesium from class II metals and copper from class III metals, the oxidation-reduction can be explained using the following formula.

If other metals are used, similar explanations can be used.

Mg(s) + Cu2+(aq) ==> Mg2+(aq) + Cu(s) + 1.5v

This reaction is an oxidation-reduction reaction in which the magnesium is oxidized to Mg2+ ions and Cu2+ is reduced to metallic copper. The result of this reaction is a transfer of electrons which provides the power required to light the flashlight lamp.

Common problems: (important)

The common problem in making a galvanic cell is that the reducing metal precipitates over the oxidizing metal and the current stops. For example when you use copper and zinc as electrodes and copper sulfate as an electrolyte, in a few seconds copper will precipitate over zinc, so the condition will change like both electrodes are copper. This condition stops the chemical reactions and production of electricity. The challenge is finding a method to prevent copper from precipitating over zinc.

One common method is using a two part container where these two parts are separated from each other using a porous material such as unglazed ceramic. (See details here) In this way one side of the container will have copper electrode and copper sulfate, and the other side will have zinc electrode and zinc sulfate.

The other method is using two separate containers and using a salt bridge to connect two containers to each other.

As shown in this picture, Fingers are used in place of a salt bridge in an electrochemical cell. The voltage produced is almost the same as the standard cell potential.

Salt Bridge is a U-shaped tube containing electrolyte, which connects two half-cells of a voltaic cell.

Salt bridge electrolyte is usually a gel containing a salt such as potassium nitrate.

To make a salt bridge, heat 150 mL of distilled water to boiling in a 400 mL beaker. Add 3 g of agar and stir the mixture as it boils until a uniform suspension forms. Remove the beaker from the heat and stir in 15 g of KNO3 until the salt dissolves. Pour the warm mixture into a U shape tube until it is completely filled. Keep it upright and let it set overnight. Once the agar is set, store the salt bridges in plastic bags to prevent drying out.

Question/ Purpose:

What do you want to find out? Write a statement that describes what you want to do. Use your observations and questions to write the statement.

The purpose of this experiment is to learn about electrochemistry and oxidation-reduction reactions through the construction and operation of a simple galvanic cell.

Find out what factors affect the voltage and the maximum electric current produced by a galvanic cell.

Identify Variables:

When you think you know what variables may be involved, think about ways to change one at a time. If you change more than one at a time, you will not know what variable is causing your observation. Sometimes variables are linked and work together to cause something. At first, try to choose variables that you think act independently of each other.

Independent variables are the types and sizes of electrodes.

Dependent variables are the voltage and the maximum electric current produced by the galvanic cell.

Controlled variables are temperature, electrolytes, and experiment procedures.


Based on your gathered information, make an educated guess about what types of things affect the system you are working with. Identifying variables is necessary before you can make a hypothesis. Following is a sample hypothesis:

Voltage between two electrodes is a function of the electrode metals and the electrolyte. If the difference between the reactivity of electrode metals is more, we expect a higher voltage.

The maximum electric current depends on the size of electrodes. As the surface of electrodes in contact with electrolyte increases, more chemical reaction will happen and as a result more electricity will be produced.

Experiment Design:

Design an experiment to test each hypothesis. Make a step-by-step list of what you will do to answer each question. This list is called an experimental procedure. For an experiment to give answers you can trust, it must have a “control.” A control is an additional experimental trial or run. It is a separate experiment, done exactly like the others. The only difference is that no experimental variables are changed. A control is a neutral “reference point” for comparison that allows you to see what changing a variable does by comparing it to not changing anything. Dependable controls are sometimes very hard to develop. They can be the hardest part of a project. Without a control you cannot be sure that changing the variable causes your observations. A series of experiments that includes a control is called a “controlled experiment.”


In this experiment an inexpensive cell is made from materials obtained locally. Test the cell by connecting it to a 1.5-volt flashlight lamp and observing whether or not the lamp will light. In addition, four or more of these cells can be connected in series to form a battery that can power a small radio or other device that operates on direct current.

See the list of material section before reading the procedure.


  1. Prepare a test lamp by soldering wire test leads to a 1.5-volt flashlight lamp. Attach alligator clips to the ends of the leads.
  2. Obtain strips of magnesium and copper, each strip should be 2.5 cm longer than the height of the beaker being used. Sand the strips until they are shiny.
  3. Tie a knot with string or use a non-reactive clamp in one end of dialysis tubing that has been soaked in distilled water. The length of tubing should be long enough so that it overlaps the edge of the beaker by 2.5 cm.
  4. Fill the dialysis tubing with the prepared copper(II) sulfate solution. Place the copper strip in this piece of dialysis tubing that is now filled with copper(II) sulfate solution and use string or a rubber band to secure the top of the dialysis tubing around the copper. Leave 2.5 cm of copper sticking out of the tubing.
  5. Fill the beaker with the prepared sodium sulfate solution, place both the dialysis tubing, containing the copper and copper(II) sulfate solution, and the magnesium strip in the beaker. Place the magnesium strip as far away as possible from the dialysis tubing. Secure both the magnesium strip and the copper strip in the dialysis tubing in the beaker by bending them over the edge of the beaker.
  6. Complete the circuit by attaching one alligator clip to the copper strip and the other to the magnesium strip.
  7. Observe and record any activity taking place at the metal strips, the solution (especially color changes), and the test lamp. Also measure and record the voltage between electrodes with and without the test lamp connected. Also measure the electric current without a test lamp in the circuit.
  8. Connect several cells in series and measure the voltage again.
  9. Try making cells with zinc or aluminum in place of the magnesium.
  10. Try making cells with larger or smaller electrodes
  11. All solutions may be flushed down the drain with water.


Dialysis tubing is a porous membrane that allows SO42- ions to pass, but blocks Zn2+ and Cu2+ ions. A cellophane tube and a ceramic tube will also work as well.

I think intestine casings (skin of sausage) may also work as a membrane, but I have not tried it. You can purchase these from your local meat market or butcher’s shop. If you try this, please let me know about the result. Send an email to info@ScienceProject.com.

The top of the membrane must not be air tight. Some hydrogen gas may form on the zinc electrode and needs to exit.

Inside the membrane sack, you can use zinc sulfate, Sodium sulfate, potassium sulfate or ammonium sulfate.

Substitute chemicals and electrodes:

As you see in the above diagram, you can substitute magnesium with zinc and sulfate salt with nitrate salt with similar results.

Avoid contact with solutions. Goggles must be worn throughout experiment.

Materials and Equipment:


0.5 M copper(II) sulfate (dissolve 125 g CuSO4ยท5H2O in distilled or deionized water and dilute to 1.00 liter)
0.5 M sodium sulfate (dissolve 71 grams of Na2SO4 in distilled or deionized water and dilute to 1.00 liter)
zinc, copper, aluminum, and magnesium strips (approximately 1 cm x 10 cm)


250-mL beaker or a mug
dialysis tubing
hook-up wire or bell wire
1.5-volt flashlight lamp with less than 100 milliamp rating
alligator clamps
crimping tool
soldering iron and resin core solder
clamps(for dialysis tubing)
sandpaper or steel wool

Modifications/Substitutions :

  1. Copper sulfate pentahydrate is available as root killer at garden supply stores. It may also be purchased from pool supliers.
  2. Aluminum, from aluminum cans, may be used if it is sanded well on both sides. Aluminum gutter nails may also be used.
  3. Copper tubing or fittings may be used in place of copper strips.
  4. Zinc can be obtained from old dry cell battery casings. This should be done carefully to avoid contact with caustic chemicals in battery. Small zinc electrodes are available at MiniScience.com and as a part of their Make Electricity Science set.
  5. Sausage casings can be used in place of dialysis tubing but diffusion is extremely rapid. Dialysis tubing is readily available in biology labs.
  6. Baby food jars or other open glass jars may be used instead of beakers.

Dialysis tubing is actually a semi-permeable membrane when used in water. This tubing usually comes in rolls and when wet, will open up into a cylindrical tube that can be tied off at the ends. The tubing is usually available at different diameters from 1 cm up to 10 cm.

Any diameter from 3 to 7 cm is good for this experiment.

Results of Experiment (Observation):

Experiments are often done in series. A series of experiments can be done by changing one variable a different amount each time. A series of experiments is made up of separate experimental “runs.” During each run you make a measurement of how much the variable affected the system under study. For each run, a different amount of change in the variable is used. This produces a different amount of response in the system. You measure this response, or record data, in a table for this purpose. This is considered “raw data” since it has not been processed or interpreted yet. When raw data gets processed mathematically, for example, it becomes results.



Summary of Results:

Summarize what happened. This can be in the form of a table of processed numerical data, or graphs. It could also be a written statement of what occurred during experiments.

It is from calculations using recorded data that tables and graphs are made. Studying tables and graphs, we can see trends that tell us how different variables cause our observations. Based on these trends, we can draw conclusions about the system under study. These conclusions help us confirm or deny our original hypothesis. Often, mathematical equations can be made from graphs. These equations allow us to predict how a change will affect the system without the need to do additional experiments. Advanced levels of experimental science rely heavily on graphical and mathematical analysis of data. At this level, science becomes even more interesting and powerful.


Using the trends in your experimental data and your experimental observations, try to answer your original questions. Is your hypothesis correct? Now is the time to pull together what happened, and assess the experiments you did.

Related Questions & Answers:

What you have learned may allow you to answer other questions. Many questions are related. Several new questions may have occurred to you while doing experiments. You may now be able to understand or verify things that you discovered when gathering information for the project. Questions lead to more questions, which lead to additional hypothesis that need to be tested.

  • What material can be used as a membrane in a voltaic cell?
  • What salts can be used as electrolyte?
  • How does the distance between electrodes affect the voltage?

Possible Errors:

  1. The metal strips used in this experiment should have a large surface area to minimize resistance within the cell. (I recommend a surface area of 15 square inches or more.)
  2. Use a voltmeter instead of the light bulb to detect small amounts of electricity.
  3. Aluminum strips may not work unless sanded thoroughly and acid treated with 6.0 M HCl.
  4. Use a lamp holder for a secure connection with light bulb and eliminating the need to do soldering.


Summerlin, L.R., and Ealy, J.L. Jr, Chemical Demonstrations-A Sourcebook for Teachers, American Chemical Society, Washington D.C, 1985, p. 115. This experiment is adapted from this source.


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