Materials Included In Kit

Additional Materials Required

Safety Precautions

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.

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. 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.

Lab Hints

• 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.

Teacher Tips

• 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)

  {12775_Tips_Equation_1}

  Reduction-half-reaction (cathode)

  {12775_Tips_Equation_2}

  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.

  {12775_Tips_Equation_3}

  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)

  {12775_Tips_Equation_1}

  Reduction-half-reaction (cathode)

  {12775_Tips_Equation_5}

  Overall reaction

  {12775_Tips_Equation_6}

  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.

Answers to Prelab Questions

1. Complete the following table summarizing the general properties of the electrodes in an electrolytic cell.
   {12775_Answers_Table_2}
2. 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.

Sample Data

Answers to Questions

1. Review your observations for the electrolysis of the silver nitrate solution.
   1. 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.

   2. 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.

2. 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.

3. 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.

Teacher Handouts

12775_Teacher1.pdf

Introduction

Electrolysis involves the use of electric current to cause an oxidation–reduction reaction to take place. This process can be used to convert metal ores to pure metals, break down water to hydrogen and oxygen gases and plate one metal on another.

Concepts

• Electrolysis
• Oxidation and reduction
• Anode and cathode

Background

Nearly all metals in nature occur as oxides, halides, sulfites, carbonates, phosphates or other molecular compounds. Although a wide variety of physical and chemical methods are used to reduce these metallic compounds, most of the processes that produce pure metals utilize electrolysis at some stage.

When an electric current is passed through a solution, chemical reactions take place where the current enters the solution and where it leaves the solution. This process is called electrolysis. A typical setup is shown in Figure 1.

The negative electrode is called the cathode and the positive electrode is called the anode. For current to flow through the solution a substance in the water must gain electrons at the cathode and another substance must lose electrons at the anode. When a substance gains electrons in a reaction, it is called a reduction reaction, and the substance is said to have been reduced. When a substance loses electrons, it is called an oxidation reaction, and the substance has been oxidized. In a metal salt solution, there are three substances in the solution; water molecules, the salt cation and the salt anion. One of these will be reduced at the cathode and one will be oxidized at the anode. If a water molecule is reduced, hydrogen gas is produced and a basic solution forms around the cathode. If the water molecule is oxidized, oxygen gas is produced and an acidic solution forms around the anode.

Experiment Overview

The purpose of this experiment is to determine the products obtained in the electrolysis of aqueous silver nitrate, copper(II) sulfate and zinc nitrate solutions. The electrolysis of salt solutions will be investigated using a “Petri dish electrolysis” setup with a 9-V battery and carbon (pencil lead) electrodes (see Figure 2).

Materials

Prelab Questions

1. Complete the following table summarizing the general properties of the electrodes in an electrolytic cell.
   {12775_PreLab_Table_1}
2. 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?

Safety Precautions

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 strong skin irritant and is 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. Wash hands thoroughly with soap and water before leaving the lab.

Procedure

Prelab Preparation

1. Attach the carbon pencil leads to the wire leads of the battery cap as follows.
2. Strip about ¼ inch of plastic coating from the end of each wire lead.
3. Place the top portion of the carbon pencil lead at the start of the bare wire at an angle (see Figure 3).
   {12775_Preparation_Figure_3}
4. Wrap the bare wire around the carbon pencil lead.
5. Take a 1-inch piece of transparent tape and carefully secure the bare wire to the carbon pencil lead, without breaking the pencil lead (see Figure 4).
   {12775_Preparation_Figure_4}
6. Repeat the steps for the remaining wire.
7. The completed battery cap and carbon pencil lead electrodes assembly is shown in Figure 5.
   {12775_Preparation_Figure_5}

Procedure

1. Place the partitioned Petri dish on a sheet of white paper towel. Label the compartments or segments of the Petri dish 1, 2 and 3 with a wax pencil or marking pen.
2. Using a 10-mL graduated cylinder, carefully pour about 6 mL of 0.2 M silver nitrate solution into the first compartment of the Petri dish until the bottom of the compartment is covered with solution. The compartment should be one-third to one-half full.
3. Connect the battery cap to the 9-V battery. Caution: Do not allow the electrodes to touch each other.
4. Using the Beral-type pipet, add 6–8 drops of bromthymol blue indicator solution to the solution. Note the initial color of the solution in the data table.
5. Hold the red (+) lead from the 9-V battery in one hand and the black (–) lead in the other hand. Keeping the electrodes as far apart as possible, dip the carbon pencil lead electrodes into the silver nitrate solution in the Petri dish.
6. Let the electric current run for 1–2 minutes while observing any changes in the silver nitrate solution.
7. Record all observations in the data table—be sure to indicate where changes take place (at the anode or the cathode). Refer to the Background section and the Prelab Questions for the properties of the electrodes.
8. Remove the carbon pencil lead electrodes from the electrolysis solution. Observe the surface of the carbon pencil leads. Record these observations in the data table. Carefully rinse the electrodes with distilled water from a wash bottle and gently pat dry on a paper towel.
9. Carefully pour about 6 mL of 0.5 M zinc nitrate solution into the second compartment of the Petri dish.
10. Repeat steps 4–8 for the electrolysis of zinc nitrate solution. Record observations in the data table.
11. Carefully pour about 6 mL of 0.2 M copper(II) sulfate solution into the third compartment of the Petri dish.
12. Repeat steps 4–8 for the electrolysis of copper(II) sulfate solution. Record observations in the data table.
13. Remove the carbon pencil lead electrodes from the wires and disconnect the battery cap from the battery.
14. Dispose of the electrolysis solutions as directed by the instructor.

Student Worksheet PDF

12775_Student1.pdf