Materials Included In Kit

Additional Materials Required

Safety Precautions

Copper(II) bromide solution is toxic by ingestion and may be irritating to the eyes, skin and the respiratory tract. Phenolphthalein is an alcohol-based solution and is a flammable liquid. Keep away from flames and heat. Sodium thiosulfate acidified solution is a body tissue irritant. The electrolysis reactions will generate small amounts of hazardous gases. Perform this experiment in a well-ventilated lab only and do not breathe the vapors. 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 lab.

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. Electrolysis of aqueous potassium iodide, sodium chloride and copper(II) bromide generates halogen–water solutions. The contents of the Petri dishes should be collected in a central waste disposal beaker located in the hood. Sodium thiosulfate solution is provided with the kit for disposal of the halogen water solutions according to Flinn Suggested Disposal Method #12a. The resulting waste solution should be allowed to sit overnight to thoroughly degas. It may then be rinsed down the drain with plenty of excess water according to Flinn Suggested Disposal Method #26b. Do not dispose of the electrolysis waste solutions directly down the drain.

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 the electrolysis of water, performed as either an experiment or a demonstration. Please see “Introduction to Electrochemistry” in Electrochemistry, Volume 17 in the Flinn ChemTopic^™ Labs series for directions.
• Students may need help correlating the color changes at the cathode in the electrolysis of potassium iodide and sodium chloride with the production of OH^– ions from the reduction of water molecules.
• The small amount of chlorine generated in the electrolysis of sodium chloride is noticeable only by a faint odor. The concentration is not strong enough to color the solution. The test for chlorine (step 11 in the Procedure) involves adding potassium iodide and starch to observe the formation of the familiar iodine–starch complex. Chlorine is a stronger oxidizing agent than iodine and therefore oxidizes iodide anions to iodine.
• See the experiment “All in the Family” in The Periodic Table, Volume 4 in the Flinn ChemTopic^™ Labs series, for a study of the reactivity and single replacement reactions of the halogens.
• Potassium iodide solution is light- and air-sensitive. Prepare the solution fresh within two weeks of its anticipated use and store the solution in a dark bottle, if possible.
• The halogen odors generated in these reactions are very faint. The halogen odors are not a hazard when the experiment is performed as written in a well-ventilated lab. Remind students, however, never to “sniff” their experiments!
• Other electrolytes, such as silver nitrate and zinc bromide, may also be used in this experiment. Both of these metal ions are more easily reduced than the hydrogen atoms in water. Electrolysis of silver nitrate generates silver metal at the cathode. Zinc bromide gives zinc metal at the cathode and bromine at the anode.
• This experiment may be “supersized” for demonstration purposes 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.
• Universal indicator may be used as the acid–base indicator in the electrolysis of potassium iodide or sodium chloride.
• The sodium thiosulfate solution is provided in the kit for safe disposal of the halogen–water electrolysis waste solutions. Place the solution in a beaker in the hood. The waste beaker may be used continuously by several class sections during the day.

Teacher Tips

• Electrolysis reactions provide a great critical-thinking exercise for students to deduce the reactions that take place. 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.
• Based on standard reduction potential values, oxidation of chloride ion to chlorine (E°_red = –1.36 V) is less favorable than oxidation of water to oxygen (E°_red = –1.23 V). However, there is a significant overvoltage for the oxidation of water, and thus chlorine is observed in the electrolysis of aqueous sodium chloride solution. The cause of the overvoltage is usually ascribed to a kinetically slow reaction at the anode. E°_red values predict the thermodynamic tendency of a reaction to occur, not how fast or slow the reaction will be.

Further Extensions

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.

Answers to Prelab Questions

1. Complete the following table summarizing the general properties of the electrodes in an electrolytic cell.
   {12873_Answers_Table_2}
2. Sodium metal is produced commercially by the electrolysis of molten sodium chloride. The byproduct of the reaction is chlorine gas. (a) Write the oxidation and reduction half-reactions for the electrolysis of molten sodium chloride. (b) Identify the substance that is oxidized and the substance that is reduced. (c) Write the balanced chemical equation for the overall reaction.
   1. Oxidation half-reaction (anode) 2Cl^–(l) → Cl₂(g) + 2e^–

      Reduction half-reaction (cathode) Na⁺(l) + e^– → Na(l)
      Note: Sodium metal is a liquid at the temperature required for the electrolysis of molten sodium chloride.

   2. Chloride anions are oxidized to chlorine gas; sodium cations are reduced to sodium metal.
   3. Overall balanced equation 2NaCl(l) → 2Na(l) + Cl₂(g)

      Note: Remind students about the need to balance electrons as well as atoms and charge when balancing the chemical equation for a redox reaction.

3. Sodium metal is easily oxidized—it is a very reactive metal. Sodium reacts spontaneously with water at room temperature to give sodium hydroxide and hydrogen gas. Would you expect to observe sodium metal in the electrolysis of aqueous sodium chloride? Explain.

   The fact that sodium is very reactive and easily oxidized suggests that it should be extremely difficult to reduce sodium cations. In aqueous sodium chloride solution, therefore, reduction of water to hydrogen gas should be more favorable than reduction of sodium cations to sodium metal. Sodium metal will not be generated in the electrolysis of aqueous sodium chloride.

Sample Data

Answers to Questions

1. The following oxidation and reduction half-reactions are possible for the electrolysis of potassium iodide solution. The solution contains water molecules, potassium ions (K⁺) and iodide ions (I^–).

   2H₂O(l) → O₂(g) + 4H⁺(aq) + 4e^–
   2H₂O(l) + 2e^– → H₂(g) + 2OH^–(aq)
   K⁺(aq) + e^– → K(s)
   2I^–(aq) → I₂(s) + 2e^–

   1. What product was formed at the anode in the electrolysis of potassium iodide solution? Explain, citing specific evidence from your observations.

      The substance formed at the anode is an oxidation product. The product is yellow, water-soluble and turns black when starch is added—iodine.

   2. What product was formed at the cathode in the electrolysis of potassium iodide solution? Explain based on your observations.

      The substance formed at the cathode is a reduction product. The product is a gas, and is accompanied by the formation of a base (phenolphthalein turned pink). The product is hydrogen, and hydroxide ions are formed as a by-product.

   3. Write the balanced chemical equation for the overall redox reaction in the electrolysis of aqueous potassium iodide. Hint: Remember to balance the electrons!

      2H₂O(l) + 2I^–(aq) → H₂(g) + I₂(aq) + 2OH^–(aq)

2. Using Question 1 as a guide: (a) Identify the products that were formed at the anode and the cathode in the electrolysis of sodium chloride solution, giving the specific evidence for their formation. (b) Write the balanced chemical equation for the overall redox reaction.
   1. The substance formed at the anode (oxidation product) is a water-soluble gas with a “swimming pool” odor—chlorine. The dark yellow color observed when potassium iodide was added is due to iodine. (Chlorine oxidizes iodide ions to iodine.) The substance formed at the cathode (reduction product) is hydrogen. Hydroxide ions are formed as a byproduct. See the observations and explanation for electrolysis of potassium iodide.
   2. 2H₂O(l) + 2Cl^–(aq) → H₂(g) + Cl₂(aq) + 2OH^–(aq)
3. Using Question 1 as a guide: (a) Identify the products that were formed at the anode and the cathode in the electrolysis of copper(II) bromide solution, giving the specific evidence for their formation. (b) Write the balanced chemical equation for the overall redox reaction.
   1. The substance formed at the anode (oxidation product) is a dark yellow, water-soluble liquid with a sharp odor—bromine. The substance formed at the cathode (reduction product) is a reddish brown solid—copper metal.
   2. Cu²⁺(aq) + 2Br^–(aq) → Cu(s) + Br₂(aq)
4. Compare the product formed at the cathode in the electrolysis of copper(II) bromide solution versus that obtained in the electrolysis of aqueous potassium iodide or sodium chloride. Explain, based on the reactivity of the metals.

   Copper metal was obtained at the cathode (reduction product) in the electrolysis of copper(II) bromide solution. This contrasts with the formation of hydrogen as the reduction product in the electrolysis of aqueous potassium iodide or sodium chloride. Copper(II) ions are therefore more easily reduced than water molecules or potassium or sodium ions The ease of reduction: Cu²⁺ > H₂O >> Na⁺, K⁺. Potassium and sodium are very reactive metals—they are easy to oxidize, their cations are difficult to reduce. Copper metal is a relatively unreactive metal—it is harder to oxidize, but its cations are easy to reduce.

5. (Optional) Consult a table of standard reduction potentials (Eº_red): Determine the minimum voltage necessary for the electrolysis of aqueous potassium iodide. Hint: Eº_cell = Eº_red (cathode) – Eº_red (anode)

   Oxidation (anode): 2I^–(aq) → I₂(s) + 2e^– Eº_red = +0.54 V
   Reduction (cathode): 2H₂O(l) + 2e^– → H₂(g) + 2OH^–(aq) Eº_red = –0.83 V
   Eº_cell = Eº_red (cathode) – Eº_red (anode) = –0.83 V – 0.54 V = –1.37 V
   The minimum cell voltage required for this nonspontaneous reaction is 1.37 V.
   Note: E° values are based on 1 M solutions of all ions, which was not the case in this experiment.

References

This kit was adapted from Electrochemistry, Flinn ChemTopic^™ Labs, Vol. 17, Cesa, I., Editor; Flinn Scientific Inc.: Batavia, IL (2005).

Introduction

Electrolysis is defined as the decomposition of a substance by means of an electric current. When an electric current is passed through water containing an electrolyte, the water molecules decompose via an oxidation–reduction reaction. Oxygen gas is generated at the anode, hydrogen gas at the cathode. The purpose of the electrolyte, such as sodium sulfate, is to provide ions that will “carry” the current 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.

Concepts

• Electrolysis
• Oxidation and reduction
• Anode and cathode
• Cell potential

Background

An electrolytic cell consists of a power source or a battery connected to two electrodes in a solution of an electrolyte. The electrodes act as external conductors and provide surfaces at which electron transfer will take place. Electrons flow from the anode, which is the site of oxidation, to the cathode, which is the site of reduction. The power source or battery serves as an electron “pump,” pushing electrons into the electrolytic cell from the negative pole and pulling electrons from the cell at the positive pole. The negative electrode, where the electrons enter the cell, is the cathode. The electrons are “consumed” in a reduction half-reaction at the cathode. Electrons are generated at the anode, the positive electrode, via an oxidation half-reaction. The migration of ions in the electrolyte solution completes the electrical circuit.

The following half-reactions occur in the electrolysis of water:

Oxidation half-reaction (anode) 2H₂O(l) → O₂(g) + 4H⁺(aq) + 4e^–
Reduction-half-reaction (cathode) 2H₂O(l) + 2e^– → H₂(g) + 2OH^–(aq)

Electrolysis of an aqueous solution may generate products other than oxygen or hydrogen if the electrolyte contains ions that are more easily oxidized or more easily reduced than water molecules. The electrolysis of aqueous silver nitrate (AgNO₃), for example, produces oxygen at the anode and silver metal at the cathode. The products of the reaction demonstrate that reduction of silver ions (Ag⁺) to silver (Ag) occurs more readily than reduction of water. The overall reaction is the sum of the oxidation and reduction half-reactions:

Oxidation half-reaction (anode) 2H₂O(l) → O₂(g) + 4H⁺(aq) + 4e^–
Reduction-half-reaction (cathode) 4Ag⁺(aq) + 4e^– → 4Ag(s)
Overall reaction 2H₂O(l) + 4Ag⁺(aq) → O₂(g) + 4Ag(s) + 4H⁺(aq)

Experiment Overview

The purpose of this experiment is to identify the products obtained in the electrolysis of aqueous potassium iodide, copper(II) bromide and sodium chloride solutions. The electrolysis reactions will be carried out in an electrolytic cell consisting of a Petri dish, a 9-V battery and carbon (pencil lead) electrodes (see Figure 1).

Materials

Prelab Questions

1. Complete the following table summarizing the general properties of the electrodes in an electrolytic cell.
   {12873_PreLab_Table_1}
2. Sodium metal is produced commercially by the electrolysis of molten sodium chloride. The by-product of the reaction is chlorine gas. (a) Write the oxidation and reduction half-reactions for the electrolysis of molten sodium chloride. (b) Identify the substance that is oxidized and the substance that is reduced. (c) Write the balanced chemical equation for the overall reaction.
3. Sodium metal is easily oxidized—it is a very reactive metal. Sodium reacts spontaneously with water at room temperature to give sodium hydroxide and hydrogen gas. Would you expect to observe sodium metal in the electrolysis of aqueous sodium chloride? Explain.

Safety Precautions

Copper(II) bromide solution is toxic by ingestion and may be irritating to the eyes, skin and respiratory tract. Phenolphthalein is an alcohol-based solution and is a flammable liquid. Keep away from flames and heat. The electrolysis reactions will generate small amounts of gases. Do not breathe the vapors. 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

1. Place the partitioned Petri dish on a sheet of white paper. Observe that the compartments or segments of the Petri dish are labeled 1, 2 and 3.
2. Carefully pour about 8 mL of 0.5 M potassium iodide solution into the first compartment of the Petri dish until the compartment is one-third to one-half full.
3. Add 3 drops of phenolphthalein solution and stir to mix.
4. Connect the battery cap to the 9-V battery. Carefully attach a “pencil lead” electrode to each alligator clip lead. Caution: Do not allow the electrodes to touch each other.
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 pencil lead electrodes into the potassium iodide solution.
6. Let the electric current run for 1–2 minutes while observing any changes in the potassium iodide solution. 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.
7. Remove the pencil lead electrodes from the electrolysis solution. Carefully rinse the electrodes with distilled water from a wash bottle and gently pat dry on a paper towel.
8. Add two drops of starch solution to the potassium iodide solution after electrolysis and record observations in the data table.
9. Carefully pour about 8 mL of 0.5 M sodium chloride solution into the second compartment of the Petri dish. Add three drops of phenolphthalein indicator solution and stir to mix.
10. Repeat steps 5–7 for the electrolysis of sodium chloride solution. Record observations in the data table.
11. After electrolysis, add 3 drops of potassium iodide solution, followed by one drop of starch, to the sodium chloride solution. Record observations in the data table.
12. Carefully pour about 8 mL of 0.2 M copper(II) bromide solution into the third compartment of the Petri dish.
13. Repeat steps 5–7 for the electrolysis of copper(II) bromide solution. Record observations in the data table.
14. Remove the pencil lead electrodes from the alligator clips and disconnect the battery cap from the battery.
15. The electrolysis products may include dilute halogen solutions (chlorine, bromine, and iodine). Working in the hood, carefully pour the contents of the Petri dish into a waste beaker containing sodium thiosulfate solution. Sodium thiosulfate will reduce the halogen waste products. Allow the beaker to stand in the hood overnight.

Student Worksheet PDF

12873_Student1.pdf