Materials Included In Kit

Additional Materials Required

Safety Precautions

To extend the life of the battery, avoid touching the positive and negative terminals to each other. Wear chemical splash goggles and chemical-resistant gloves. 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. The electrolysis solution may be disposed of down the drain with plenty of excess water according to Flinn Suggested Disposal Method #26b.

Lab Hints

• Enough materials are provided in this Super Value Kit for 5 classes of 30 students each, working in pairs (75 total student groups). The Petri dishes and battery caps with wire leads are completely reusable.
• The laboratory work for this experiment can easily be completed in 20–30 minutes. The remaining time in a typical 50-minute lab period may be used to carry out some complementary demonstrations. See the Teaching Tips.
• Sodium sulfate is used as a source of dissolved ions to increase the current flow through the solution. In the absence of added electrolyte, no reaction will occur when the battery is connected to the pencil leads—there are no ions to “carry” the current through the solution. The rate of electrolysis increases as the concentration of sodium sulfate increases (compare 0.1 M and 1 M solutions). The conductivity of a sodium sulfate solution versus pure water may be demonstrated using a conductivity tester (available from Flinn Scientific, Catalog No. AP5355).
• The sulfate ion is an extremely weak base (the pK_a for its conjugate acid, HSO₄^–, is 2.0). The initial indicator color for the electrolysis solution may be more blue-green rather than green. According to the Merck Index, the pH of a sodium sulfate solution is 6.0–7.6. Test a small amount of the sodium sulfate stock solution with bromthymol blue indicator before class—the solution should turn green. If the solution is blue, add one drop of 1 M hydrochloric acid to the stock solution. If the solution is yellow, add one drop of 1 M sodium hydroxide to the stock solution.
• The sodium sulfate electrolysis solution may be reused by several classes during the day. Recycling the solution in this way means the indicator will already be present—instruct students to omit step 3 in the Procedure. Adjust the pH of the electrolysis solution as described above, if necessary.
• The odor of chlorine may be observed at the anode if sodium chloride is substituted for sodium sulfate as the electrolyte in the electrolysis solution. Based on their standard reduction potentials, oxidation of chloride ion to chlorine (Eº = –1.36 V) is less favorable than oxidation of water to oxygen (Eº = –1.23 V). However, there is a significant overvoltage for oxidation of water, and oxidation of chloride competes with oxidation of water under typical electrolysis conditions. Although the cause of the overvoltage is poorly understood, it is generally believed to be due to a kinetically slow reaction at the anode.
• Students may need to bend the wires to properly position the carbon pencil leads in step 4.

Teacher Tips

• A large, demonstration-scale version of this experiment uses the Hoffman electrolysis apparatus. Using the Hoffman apparatus makes it possible to measure the volume of gas generated at each electrode, collect the gases, and test their properties. If the time and current are measured, the amount of hydrogen gas collected may also be used for quantitative calculations of the Faraday constant.
• In the Flinn kit, AP6374, “Micro Mole Rockets—Hydrogen and Oxygen Ratio,” the gas mixture generated in the microscale electrolysis of water is collected in a jumbo pipet bulb. Igniting the gas mixture with a piezoelectric igniter propels the resulting “bulb rocket” across the room—a great way to compare the properties of a 2:1 stoichiometric mixture of hydrogen and oxygen versus pure hydrogen.
• 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.
• See the following website for information about the principles and design of fuel cells. http://www.netl.doe.gov/coolscience/teacher/lesson-plans/lesson6.pdf (accessed March 2005).

Answers to Prelab Questions

1. A decomposition reaction may be defined as any reaction in which one reactant, a compound, breaks down to give two or more products. Write the balanced chemical equation for the decomposition of water to its elements.
   {12090_Answers_Equation_1}
2. 1. Assign oxidation states to the hydrogen and oxygen atoms in each substance in the above chemical equation.

      Oxidation state of H in H₂O = +1. Oxidation state of O atom in H₂O = –2.
      Oxidation state of H and O in elemental H₂ and O₂ = 0.

   2. Based on the changes in oxidation states for each atom, identify the atom that is oxidized and the atom that is reduced in the decomposition of water.

      The oxygen atom is oxidized (oxidation state increases from –2 to 0). The hydrogen atom is reduced (oxidation state decreases from +1 to 0).

3. Balance the following oxidation and reduction half-reactions for the decomposition of water. Hint: Hydrogen ions (H⁺) and hydroxide ions (OH^–) are required to balance atoms and charge.

   2 H₂O → O₂ + 4 H⁺ + 4 e^–
   2 H₂O + 2 e^– → H₂ + 2 OH^–

4. Explain how the oxidation and reduction half-reactions may be combined to give the balanced chemical equation for the decomposition of water. What happens to the electrons and to the H⁺ and OH^– ions?

   The electrons must balance or “cancel out” when the oxidation and reduction half-reactions are combined. The reduction half-reaction must therefore be multiplied by a factor of two. H⁺ and OH^– ions generated in the individual half-reactions combine to form water molecules.

Sample Data

Answers to Questions

1. Suggest an explanation for the initial indicator color of the electrolysis solution.

   The initial indicator color of the electrolysis solution was green—the solution is near neutral, pH = 6.0–7.6. This is consistent with the fact that sodium sulfate is a neutral salt (like sodium chloride). The resulting solution is neither acidic nor basic.

2. Describe at least three observations that indicate a chemical reaction has occurred during the electrolysis of water.

   Signs of a chemical reaction: Formation of gas bubbles at each electrode; indicator color change to yellow at the positive electrode; indicator color change to blue at negative electrode. Note: The gas bubbles appear visibly different at the cathode versus the anode.

3. What are the two functions of the pencil lead electrodes?

   The electrodes act as external conductors for the electric current between the battery and the solution and also provide a surface for the chemical reactions.

4. 1. Compare the color changes observed at the positive (+) and negative (–) electrodes. What ions were produced at each electrode?

      The indicator color changed to yellow at the (+) electrode. This is due to the formation of H⁺ (H₃O⁺) ions. Bromthymol blue is yellow in acidic solutions (pH <6), when the concentration of H⁺ ions is greater than the concentration of OH^– ions. the indicator color changed to blue at the negative electrode. this is due to the formation of OH^– ions. bromthymol blue is blue in basic solutions (ph>7.6), when the concentration of OH^– ions is greater than the concentration of H⁺ ions.

   2. Write out the oxidation and reduction half-reactions for the decomposition of water and identify which reaction occurred at each electrode, based on the indicator color changes.

      2H₂O(l) → 4H⁺(aq) + O₂(g) + 4e^– Oxidation occurred at the (+) electrode. (Production of H⁺ ions)
      2H₂O(l) + 2e^– → H₂(g) + 2OH^–(aq) Reduction occurred at the (–) electrode. (Production of OH^– ions)

5. Compare the rates of gas evolution at the positive (+) and negative (–) electrodes. What gas was produced at each electrode? Explain, based on the balanced chemical equation for the decomposition of water (see Prelab Question 1).

   The rate of gas evolution was greater at the negative electrode, where hydrogen gas is formed. According to the balanced chemical equation for the decomposition of water, two moles of hydrogen gas are formed for every mole of oxygen gas that is released.

6. Suggest an explanation for the final indicator color of the mixed electrolysis solution.

   The final indicator color of the mixed electrolysis solution was green (neutral)—the total number of H⁺ ions produced at the positive electrode is equal to the total number of OH^– ions produced at the negative electrode. The balanced chemical equation for the overall reaction at both electrodes shows there is no net excess of either ion. (The ions combine to produce water molecules.)

7. Think about the flow of electrons and current in the electrolysis of water. What do the positive and negative signs on a battery signify?

   Electrons flow from the negative battery terminal to the negative electrode, where they are “consumed” in the reduction half-reaction. Oxidation occurs at the positive electrode, where electrons are released and flow into the positive terminal on the battery. The battery acts as an “electron pump,” pushing electrons into one electrode and pulling them from the other electrode. Electric current flows through the electrolysis solution via the migration of ions.

   {12090_Answers_Figure_4}

   Note: In all types of electrochemical cells, electrons carry the current through the external wire, while ions carry the current through the solution. Anions move toward the anode, cations move toward the cathode.

8. (Optional) Decomposing water to its elements requires energy in the form of electricity. The reverse process, combining hydrogen and oxygen to form water, may be used to generate electricity in a fuel cell. Research and describe the basic features of a fuel cell.

   A fuel cell is a device for the direct production of electricity from the energy released in a chemical reaction. Although in theory many different fuels may be used as the energy source, most of the research and engineering today refers to hydrogen fuel cells, in which hydrogen is the energy source. The chemical reaction is the combustion (combination) reaction of hydrogen with oxygen to produce water and energy.

References

This experiment has been adapted from Flinn ChemTopic^™ Labs, Volume 17, Electrochemistry; Cesa, I., Ed., Flinn Scientific: Batavia, IL, 2005.

Introduction

Electrochemistry is the study of the relationship between electrical forces and chemical reactions. There are two basic types of electrochemical processes. In a voltaic cell, commonly known as a battery, the chemical energy from a spontaneous oxidation–reduction reaction is converted into electrical energy. In an electrolytic cell, electricity from an external source is used to “force” a nonspontaneous chemical reaction to occur. What chemical reaction will take place when an electric current flows through water?

Concepts

• Electrochemistry
• Electrolysis
• Oxidation–reduction
• Anode vs. cathode

Background

The first electrochemical process to produce electricity was described in 1800 by the Italian scientist Alessandro Volta, a former high school teacher. Acting on the hypothesis that two dissimilar metals could serve as a source of electricity, Volta constructed a stacked pile of alternating silver and zinc plates separated by pads of absorbent material soaked in saltwater. When Volta moistened his fingers and repeatedly touched the top and bottom metal plates, he experienced a series of small electric shocks. The “voltaic pile,” as it came to be called, was the first battery—a chemical method of generating an electric current. Within months, William Nicholson and Anthony Carlisle in England attempted to confirm the production of electric charges on the upper and lower plates in a voltaic pile using an electroscope. In order to connect the plates to the electroscope, Nicholson and Carlisle added some water to the uppermost metal plate and inserted a wire to the electroscope. To their surprise, Nicholson and Carlisle observed the formation of a gas, which they identified as hydrogen. Nicholson and Carlisle then filled a small tube with river water and inserted wires from the voltaic pile into each end of the tube. Two different gases were generated, one at each wire—Nicholson and Carlisle had discovered electrolysis.

Experiment Overview

The purpose of this experiment is to investigate the electrolysis of water in an electrochemical cell. Two carbon pencil “leads” will be inserted into the opposite ends of a Petri dish containing water, sodium sulfate, and bromthymol blue. An electric current will be passed through the solution by connecting the pencil leads to the positive and negative terminals of a 9-volt battery (see Figure 1 in the Prelab Questions). The pencil leads act as external conductors and provide a surface for the chemical reaction. Sodium sulfate, an ionic compound, is needed to improve the current flow through the solution. Bromthymol blue, an acid–base indicator, will help to identify the changes occurring in the solution as the electrolysis proceeds. Bromthymol blue is yellow in acidic solutions (pH <6.0), blue in basic solution (ph>7.6), and various shades of green at intermediate pH values (pH = 6.0–7.6).

Materials

Prelab Questions

Recall that any oxidation–reduction reaction may be written as the sum of two half-reactions, an oxidation half-reaction and a reduction half-reaction. Electrons flow from the substance that is oxidized (which loses electrons), to the substance that is reduced (which gains electrons). If the half-reactions are separated, the electrons will flow through an external conductor rather than through the solution. This is the basis of electrochemistry. In electrolysis, the electron flow is not spontaneous, but rather is “forced” by a battery.

1. A decomposition reaction may be defined as any reaction in which one reactant, a compound, breaks down to give two or more products. Write the balanced chemical equation for the decomposition of water to its elements.
2. 1. Assign oxidation states to the hydrogen and oxygen atoms in each substance in the above chemical equation.
   2. Based on the changes in oxidation states for each atom, identify the atom that is oxidized and the atom that is reduced in the decomposition of water.
3. Balance the following oxidation and reduction half-reactions for the decomposition of water. Hint: Hydrogen ions (H⁺) and hydroxide ions (OH^–) are required to balance atoms and charge.

   H₂O → O₂ + H⁺ + e^–
   H₂O + e^– → H₂ + OH^–

4. Explain how the oxidation and reduction half-reactions may be combined to give the balanced chemical equation for the decomposition of water. What happens to the electrons and to the H⁺ and OH^– ions?
   {12090_PreLab_Figure_1_Petri dish setup for the electrolysis of water}

Safety Precautions

To extend the life of the battery, avoid connecting the positive and negative terminals to each other. Wear chemical splash goggles, chemical-resistant gloves and a chemical-resistant apron. Wash hands thoroughly with soap and water before leaving the lab.

Procedure

Preparation

1. Attach the pencil leads to the wire leads of the battery clip as follows.
2. Strip about ¼ inch of plastic coating from the end of each wire lead.
3. Place the top portion of the pencil lead at the start of the bare wire on an angle (see Figure 2).
   {12090_Preparation_Figure_2}
4. Wrap the bare wire around the pencil lead.
5. Take a 1" piece of transparent tape and carefully secure the bare wire to the pencil lead without breaking the lead (see Figure 3).
   {12090_Preparation_Figure_3}
6. Repeat these steps for the the remaining wire.

1. Attach the battery cap with the carbon leads to the 9-V battery. The battery may be placed in a support clamp, if needed, to prevent tension when the carbon pencil leads are placed into the solution (step 4).
2. Obtain about 40 mL of 0.5 M sodium sulfate solution in a small beaker. Using a Beral-type pipet, add about 3 mL of bromthymol blue indicator and stir the solution to evenly distribute the indicator color. Observe and record the initial indicator color of the solution.
3. Carefully pour the electrolysis solution into the Petri dish until the liquid level fills about two-thirds of the dish.
4. Place the pencil leads at opposite sides of the Petri dish. Be sure the leads extend into the solution (see Figure 1 in Prelab Questions).
5. Observe and record all changes as the current flows through the electrolysis solution. Be specific—compare the changes at the pencil lead electrodes attached to the positive (+) and negative (–) terminals of the battery.
6. Allow the current to flow through the solution for about 5 minutes.
7. Remove the pencil leads from the solution.
8. Disconnect the battery cap from the battery and return both to their proper location.
9. Pour the solution from the Petri dish into a small beaker. Observe and record the final indicator color of the mixed solution.
10. Dispose of the final solution as directed by the instructor.

Student Worksheet PDF

12090_Student1.pdf