Electrolysis of Metal Salts
Student Laboratory Kit
Materials Included In Kit
Bromthymol blue indicator solution, 0.04%, 50 mL
Copper(II) sulfate solution, CuSO4, 0.5 M, 100 mL
Silver nitrate solution, AgNO3, 0.2 M, 100 mL
Zinc nitrate solution, Zn(NO3)2, 0.5 M, 100 mL
Battery cap with wire leads, 15
Pencil lead electrodes, 0.9-mm, 30
Petri dishes, partitioned, 3-way, 15
Additional Materials Required
Graduated cylinders, 10-mL
Wax pencil or marking pen
Copper(II) sulfate solution is slightly toxic by ingestion. Silver nitrate solution is corrosive and will stain skin and clothing. Zinc nitrate solution is a skin irritant and moderately toxic by ingestion. Avoid contact of all chemicals with eyes and skin. Wear chemical splash goggles, chemical-resistant gloves and a chemical-resistant apron. Please review current Safety Data Sheets for additional safety, handling and disposal information. Remind students to wash their hands thoroughly with soap and water before leaving the laboratory.
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 electrolysis waste solutions, the zinc chloride solution, and the copper(II) bromide solution may be disposed of according to Flinn Suggested Disposal Method #26b. The silver nitrate solution may be disposed of according to Flinn Suggested Disposal Method #11.
- The laboratory work for this experiment can easily be completed in a typical 50-minute lab period. The experiment works best as a follow-up to an experiment or demonstration of the electrolysis of water.
- Students may need to bend the wires to properly position the carbon leads in step 3.
- This experiment may be “supersized” by carrying out the reactions in U-tubes with carbon rod electrodes and a 6-V lantern battery as the power source. About 30–50 mL of electrolyte solution will be needed, depending on the size of the U-tubes.
- The Supplementary Information in the Teacher PDF contains instructions for building a combined battery–electrode assembly using two #2 pencils and a 9-V battery with snaps.
- If your lab schedule does not allow you to have students perform this activity as an experiment, do it as a demonstration instead. No procedural changes are necessary—simply place the partitioned Petri dish on an overhead projector and follow the instructions for a great critical-thinking exercise. Write down all of the possible oxidation and reduction half-reactions for each salt on the board, and then have students identify the actual products based on their observations.
- Have individual student groups research and then present a class seminar on (a) the historical role of electrolysis in the discovery of potassium, sodium, magnesium, calcium, strontium and barium; or (b) the modern importance of electrolysis in the production of industrial chemicals, including aluminum, sodium hydroxide, chlorine, etc.
- The discovery of current electricity by Alessandro Volta in 1800 led to the almost immediate discovery of electrolysis, which led, in turn, to the rapid discovery of new chemical elements. Humphry Davy, a professor at the Royal Institution in London, began extensive studies in electrochemistry that culminated in 1807–1808 with his discoveries of the metals potassium, sodium, magnesium, calcium, strontium and barium.
- Before the late 1880s, aluminum was considered a precious metal. The process for converting the very abundant ore bauxite, composed mostly of aluminum oxide, to aluminum used sodium and was very expensive. In 1886, two men, Charles Martin Hall (1863–1914) and Paul Heroult (1863–1914), independently developed an electrolytic process to produce affordable pure aluminum. The process, called the Hall–Heroult process, starts with a molten solution of aluminum oxide, Al2O3, and the aluminum containing salt cryolite, Na3ALF6. It is this molten solution that is electrolyzed to produce aluminum metal.
- A more detailed background section is included.
The following half-reactions occur in the electrolysis of water.
Oxidation half-reaction (anode)
When an electric current is passed through an aqueous solution of sodium sulfate, the water molecules decompose via an oxidation–reduction reaction. Oxygen gas is generated at the anode, hydrogen gas at the cathode. The sodium sulfate acts as an electrolyte, increasing the current flow through the solution. Depending on the nature of the electrolyte, different reactions may take place at the anode and the cathode during the electrolysis of an aqueous solution.
For an oxidation–reduction to be spontaneous, the cell potential, E°cell, for the reaction must be greater than zero. E°cell for the decomposition of water is equal to the cathode potential, E°cathode, minus the anode potential, E°anode.
E°cell = (–0.83 – 1.23)V = – 2.06 V
For the reaction to occur, a voltage of at least 2.06 volts must be applied.
The electrolysis of a metal salt solution may generate pure metals rather than hydrogen if the metal ions are more easily reduced than water molecules. The electrolysis of aqueous tin(II) nitrate, Sn(NO3)2, for example, generates oxygen at the anode and tin metal at the cathode. The products of the reaction demonstrate that reduction of tin(II) ions (Sn2+) to tin (Sn) occurs more readily than reduction of water. The overall reaction is the sum of the oxidation and reduction half-reactions.
Oxidation half-reaction (anode)
A solution of aluminum nitrate, for example, will not have aluminum ions reduced to aluminum metal during electrolysis. Aluminum has a reduction potential of –1.66 volts, which is lower than that of water. Water will always be reduced at the cathode, and aluminum will remain in solution.
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
Engaging in argument from evidence
Disciplinary Core Ideas
MS-PS1.A: Structure and Properties of Matter
MS-PS1.B: Chemical Reactions
MS-PS2.B: Types of Interactions
HS-PS1.A: Structure and Properties of Matter
HS-PS1.B: Chemical Reactions
HS-PS2.B: Types of Interactions
Cause and effect
Scale, proportion, and quantity
Energy and matter
Structure and function
Stability and change
MS-PS1-2. Analyze and interpret data on the properties of substances before and after the substances interact to determine if a chemical reaction has occurred.
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.
Answers to Prelab Questions
- Complete the following table summarizing the general properties of the electrodes in an electrolytic cell.
- The pH indicator bromthymol blue is added to a solution of zinc chloride. Bromthymol blue is yellow in an acidic solution, blue in a basic solution and green in a neutral solution. A current is passed through a solution of zinc sulfate. A silver substance plates out on the surface of the cathode. The solution around the anode turns yellow. What has been oxidized and what has been reduced?
The zinc cation has been reduced to zinc metal, which plated out on the cathode. At the anode, water was oxidized, producing oxygen and an acidic solution.
Answers to Questions
- Review your observations for the electrolysis of the silver nitrate solution.
- What product was formed at the anode in the electrolysis of silver nitrate solution? Explain, citing specific evidence from your observations.
Gas bubbles were produced at the anode. The yellow color indicates an acidic solution was produced around the anode. Oxygen was produced at the anode.
- What product was formed at the cathode in the electrolysis of silver nitrate solution? Explain based on your observations.
Silver metal produced at the electrode. Silver solid covered the cathode and trailed into the solution. No bubbles formed, eliminating hydrogen as a reduction product.
- Using Question 1 as a guide: Identify the products that were formed at the anode and the cathode in the electrolysis of zinc nitrate solution, giving the specific evidence for their formation.
Oxygen was produced at the anode. Bubbles formed on the electrode and the solution turned a deeper yellow, indicating a more acidic solution. The solid on the cathode electrode was gray and no bubbling occurred. The solid was zinc.
- Using Question 1 as a guide: Identify the products that were formed at the anode and the cathode in the electrolysis of copper(II) sulfate solution, giving the specific evidence for their formation.
Oxygen was probably the product at the anode. Bubbles were produced, but no determination of the pH of the solution could be made. Copper was produced at the cathode.
||Blueprint T-Shirts—Student Laboratory Kit
||U-Shaped Drying Tube, 100 mm
||Water, Distilled, 4 L
||Batteries, Transistor Battery (Alkaline), 9 V
||Cylinder, Borosilicate Glass, 10 mL
||Bottles, Washing, Polyethylene, 500-mL
||Wax Pencil Set, Heat Resistant