Materials Included In Kit

Additional Materials Required

Prelab Preparation

• Iron(II) ammonium sulfate solution, 0.1 M: Prepare fresh within 1–2 days of use. Dissolve 3.9 g of iron(II) ammonium sulfate hexahydrate [Fe(NH₄)₂(SO₄)₂•6H₂O] in 50 mL of distilled or deionized water. Stir to dissolve and dilute to 100 mL with water. Note: Iron(II) solutions will slowly oxidize in air. This reagent is also called ferrous ammonium sulfate.
• Vitamin C solution, 0.2%: Prepare fresh before use. Dissolve 100 mg of ascorbic acid in about 25 mL of distilled or deionized water. Stir to dissolve and dilute to 50 mL with water. Commercial Vitamin C tablets may be used as a source of ascorbic acid—dissolve a 500-mg tablet in 250 mL of water. Filter the solution to remove any insoluble binder residue.

Safety Precautions

Hydrochloric acid is a corrosive liquid. It causes severe skin burns and eye damage and is toxic by ingestion or inhalation. Sodium hypochlorite solution is a corrosive liquid and moderately toxic by ingestion and inhalation. The solution reacts with concentrated acids to generate poisonous chlorine gas. Potassium ferricyanide and potassium thiocyanate solutions are irritating to skin and eyes and may evolve poisonous fumes upon heating or in contact with concentrated acids. Iron(III) chloride and sodium sulfite solutions are slightly toxic by ingestion and may be irritating to skin and eyes. Avoid contact of all chemicals with eyes and skin. Wear chemical splash goggles, chemical-resistant gloves and a lab coat or chemical-resistant apron.

Please note the following hazards for instructor preparation of the dilute solutions described above from the solids. Potassium permanganate is an oxidizing solid and may intensify a fire. Keep away from heat, sparks and open flames. Iron(III) chloride is slightly toxic by ingestion and may cause skin irritation and serious eye damage. Potassium thiocyanate is toxic by ingestion, inhalation and skin absorption. Avoid breathing dust or fumes and do not heat the solid. Potassium ferricyanide is slightly toxic by ingestion and causes skin and eye irritation. Do not heat. Because of the potential for hazardous reactions between some of the reagents used in this experiment, remind students to follow directions carefully and not to perform unauthorized experiments. 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. Potassium permanganate, sodium iodide, and Vitamin C solutions have short shelf lives. The contents of the reaction plates and excess sodium iodide and Vitamin C solutions may be rinsed down the drain with water according to Flinn Suggested Disposal Method #26b. Excess potassium permanganate solution may be reduced with sodium thiosulfate according to Flinn Suggested Disposal Method #12a. Save the remaining solutions in properly labeled bottles for future use.

Lab Hints

• The laboratory work for this microscale experiment can easily be completed within a typical 2-hour lab period. Review the Prelab Assignment prior to lab.
• Iron(II) ammonium sulfate, also known as ferrous ammonium sulfate or Mohr’s salt, is more stable than iron(II) sulfate and is commonly used to prepare standard solutions of iron(II) ions. Solutions of iron(II) ions are air- and light-sensitive—Fe²⁺ ions slowly oxidize in air, especially in the presence of acids or bases. For best results, prepare the iron(II) ammonium sulfate solution fresh each day. The solution should be pale green. Iron(II) sulfate solutions will turn noticeably yellow in the reaction plate over the course of the experiment.
• To reduce the number of disposable pipets or eyedroppers required by multiple lab sections, place solutions in labeled pipet sets or dropper bottles. When the lab is over, empty and rinse the pipets or dropper bottles and save them for use in subsequent years.
• In Part A, the dark red color due to the iron(III)–thiocyanate complex ion will slowly fade due to oxidation of thiocyanate ions by excess hydrogen peroxide (well B1).
• Some of the solutions will stain the reaction plates if left in the wells too long. The reaction plates should be rinsed well with distilled water prior to use. Remind students to rinse the wells immediately after finishing the procedure. Stains due to iron(III) compounds, potassium permanganate and Prussian blue may be removed by rinsing the wells with a dilute sodium thiosulfate solution.
• Other oxidizing and reducing agents that may be tested in this experiment include dilute nitric acid, bromine water, sodium thiosulfate, sodium nitrite and oxalic acid. Use 0.1 M solutions.
• The results obtained with sodium iodide and sodium bromide in Part B may be used to estimate the standard reduction potential for the reduction of iron(III) to iron(II). The cell potential for a redox reaction (E°_cell) is equal to the difference between the reduction potential for the reduction half-reaction (E°_red) and the reduction potential for the oxidation halfreaction (E°_ox). For a spontaneous reaction, E°_cell must be greater than zero [E°_cell = E°_red – E°_ox > 0]. The reduction potential for the I₂/I^– half-reaction is 0.54 V while the reduction potential for the Br₂/Br^– half-reaction is 1.08 V. From the observed reaction of iodide ions with Fe³⁺, we conclude that E°_red – 0.54 > 0, or E°_red > 0.54 V. Bromide ions do not react with Fe³+: E°_red – 1.08 < 0, or E°_red < 1.08 V. The standard reduction potential for the Fe³⁺/Fe²⁺ half-reaction is in the range 0.54 V < E°_red < 1.08 V. The literature value is 0.77 V (for 1 M solutions at 25 °C).
• Although all oxidation–reduction reactions can be analyzed in terms of electron transfer, it is misleading in many cases to say that oxidation and reduction actually take place via an electron transfer mechanism. There are three official IUPAC definitions for oxidation: (1) complete or net removal of one or more electrons; (2) increase in oxidation state of an atom within a compound; and (3) gain of oxygen and/or loss of hydrogen.
• Interconversion of iron(II) and iron(III) ions is important in the absorption, storage and utilization of iron by the body. Iron(II) compounds are more easily absorbed by the body than iron(III) compounds. Iron is stored in the body, however, in its oxidized iron(III) form as part of a protein called ferritin. Iron is also an essential component of many enzymes that catalyse oxidation–reduction reactions in the body. The iron atoms in these enzymes reversibly alternate between the +2 and +3 oxidation states.
• Iron occurs in two forms in foods—heme iron, which is found in meat, poultry and fish, and nonheme iron, which comes from plant sources. Heme iron is easily absorbed by the body and is the most significant source of iron. The rate of absorption of nonheme iron is much slower than that of heme iron and is strongly influenced by Vitamin C and other dietary factors.

Answers to Prelab Questions

1. Identify the oxidation state of the chlorine atom in each substance shown.
   {14039_PreLabAnswers_Table_1}
2. Potassium iodate (KIO₃) is a strong oxidizing agent and will oxidize Fe²⁺ ions to Fe³⁺. In doing so, iodate ion (IO₃^–) is reduced to elemental iodine (I₂).
   1. Use the oxidation state rules (see the Background section) to assign oxidation states to the iodine atoms in iodate ion (IO₃^–) and iodine (I₂).

      For iodate: According to rule 5, the oxidation state of each oxygen atom is –2.

      According to rule 7: (3 x –2) + (Oxidation state of I) = –1.
      The oxidation state of the iodine atom in IO₃^– is +5.

      For iodine: According to rule 1, the oxidation state of each iodine atom in I₂ is zero.

   2. The half-reaction for the reduction of iodate is shown. Use the difference in oxidation states for the iodine atoms in IO₃^– and I₂ to determine the number of electrons gained in this half-reaction. Hint: Hydrogen ions (H⁺) and water molecules (H₂O) are required to balance mass and charge.

      The difference in oxidation state for the iodine atom in iodate ion and iodine is (+5 – 0) = 5. Each iodate ion gains five electrons as the oxidation state of the iodine atom is reduced from +5 to zero.

      {14039_PreLabAnswers_Reaction_1}

      Note: The use of the fractional coefficient ½ for iodine in the balanced equation makes it easier to identify the number of electrons gained by each iodine atom. The equation may be multiplied by two to obtain whole-number coefficients for each substance. Use a “number line” to help students count the number of electrons corresponding to the change in oxidation state.

      {14039_PreLabAnswers_Figure_2}
3. Combine the oxidation half-reaction for Fe²⁺ (see the Background section) with the reduction half-reaction for iodate (Question 2b) and write the balanced equation for the overall redox reaction of Fe²⁺ with IO₃^–. Hint: The number of electrons on each side must cancel out.

   Oxidation half-reaction: Fe²⁺ → Fe³⁺ + e^–
   The oxidation half-reaction must be multiplied by a factor of five to balance the number of electrons lost by iron(II) with the number of electrons gained by one iodate ion.

   {14039_PreLabAnswers_Reaction_2}

Sample Data

Laboratory Report

Table A. Reactions of Iron(II) Ions with Oxidizing Agents

Table B. Reactions of Iron(III) Ions with Reducing Agents

Answers to Questions

1. How can potassium thiocyanate be used to confirm that Fe²⁺ ions have been oxidized to Fe³⁺ in Part A?

   A solution of Fe³⁺ ions turns dark red when potassium thiocyanate is added. If a test mixture in Part A turns red when KSCN is added, then Fe²⁺ ions have been oxidized to Fe³⁺ ions.

2. Use the oxidation state rules to assign oxidation states for the indicated atoms in each oxidizing agent and its product (Part A).
   {14039_Answers_Table_4}
3. Fill in the blanks to show the number of electrons involved in each half-reaction for the oxidizing agents identified in Question 2.
   {14039_Answers_Equation_6}
4. Combine the oxidation half-reaction for Fe²⁺ (see the Background section) with the appropriate half-reaction from Question 3 and write the balanced equation for the overall redox reaction of Fe²⁺ with (a) permanganate ion, (b) hydrogen peroxide and (c) hypochlorite ion.
   1. 5Fe²⁺(aq) + MnO₄^–(aq) + 8H⁺(aq) → 5Fe³⁺(aq) + Mn²⁺(aq) + 4H₂O(l)
   2. 2Fe²⁺(aq) + H₂O₂(aq) + 2H⁺(aq) → 2Fe³⁺(aq) + 2H₂O(l)
   3. 2Fe²⁺(aq) + OCl^–(aq) + H₂O(l) → 2Fe³⁺(aq) + Cl^–(aq) + 2OH^–(aq)
5. An oxidizing agent is a substance that causes the oxidation of another reactant in a redox reaction. The oxidation state of the oxidizing agent decreases and the oxidizing agent itself undergoes reduction during the reaction.
6. How can potassium ferricyanide be used to confirm that Fe³⁺ ions have been reduced to Fe²⁺ in Part B?

   A solution of Fe²⁺ ions turns dark blue when potassium ferricyanide is added. If a test mixture in Part B turns blue (or blue-green) when K₃Fe(CN)₆ is added, then Fe³⁺ ions have been reduced to Fe²⁺ ions.

7. 1. Sulfite ion (SO₃^2–) is a strong reducing agent. Assign oxidation states to the sulfur atom in SO₃^2– and its product, sulfate ion (SO₄^2–).

      For sulfite: [(3 x –2) + (Ox. State S) = –2]. Oxidation state of sulfur = +4.
      For sulfate: [(4 x –2) + (Ox. State S) = –2]. Oxidation state of sulfur = +6.

   2. Fill in the blank to show the number of electrons in the following half-reaction.
      {14039_Answers_Equation_7}
      The difference in oxidation states for the sulfur atom is +2. Each sulfite ion loses two electrons when it is oxidized to sulfate.
   3. Write the balanced equation for the overall redox reaction of Fe³⁺ with sulfite ion.

      2Fe³⁺(aq) + SO₃^2–(aq) + H₂O(l) → 2Fe²⁺(aq) + SO₄^2–(aq) + 2H⁺(aq)

8. A reducing agent is a substance that causes the reduction of another substance in a redox reaction. The oxidation state of the reducing agent increases and the reducing agent itself undergoes oxidation during the reaction.
9. Based on the observations in Part B, which halide—bromide ion or iodide ion—is the stronger reducing agent? Explain your reasoning.

   Sodium iodide reduced Fe³⁺ ions, whereas sodium bromide did not. Therefore, iodide ion is a stronger reducing agent than bromide ion.

10. Iron(II) compounds in foods are more easily absorbed by the body than iron(III) compounds. Vitamin C improves the absorption of dietary iron. Explain based on your observations in this experiment.

    Vitamin C is a good reducing agent and will keep the iron in its reduced iron(II) form. Note: Vitamin C is called an antioxidant. An antioxidant prevents the oxidation of biological molecules in cells and cell membranes. Oxygen and ozone in the atmosphere make the air we breathe a strongly oxidizing environment. Oxidizing agents such as hydrogen peroxide and nitric oxide are normal byproducts of cellular metabolism and cell processes. Vitamin C reduces and scavenges both internally and externally produced oxidants before they can react with important biological molecules.

11. Suggest a possible reason for the results obtained using pineapple juice in this experiment.

    Pineapple juice acted as a reducing agent, reducing Fe³⁺ ions to Fe²⁺. Pineapple juice is a natural source of Vitamin C, a strong reducing agent.

12. Which of the following reactions are not redox reactions?
    {14039_Answers_Equation_8}

    Reactions b and c are not redox reactions.

Introduction

Iron exists in the body in two forms—iron(II), Fe²⁺, and iron(III), Fe³⁺, ions and their compounds. Interconversion of the two forms of iron takes place via the loss or gain of an electron. Investigate the role of electron transfer in the reactions of iron(II) and iron(III) compounds with oxidizing and reducing agents, respectively.

Concepts

• Oxidation–reduction
• Oxidation state
• Half-reactions
• Oxidizing and reducing agents

Background

Oxidation–reduction reactions are a major class of chemical reactions. An oxidation–reduction, or redox, reaction is defined as any reaction in which electrons are transferred from one substance to another. Oxidation occurs when a substance loses electrons. Because any electrons lost by one reactant must be transferred to another reactant, oxidation and reduction always occur together. Reduction occurs when a substance gains electrons.

Substances that are used to cause the oxidation or reduction of another substance are called oxidizing and reducing agents, respectively. The substance that accepts electrons in a redox reaction is called the oxidizing agent—by accepting electrons, it is itself reduced but it causes the oxidation of another substance. Similarly, the substance that loses electrons in a redox reaction is called the reducing agent because it causes the reduction of another substance.

Combination reactions of elements in their neutral ground states to form binary ionic compounds may also be classified as redox reactions. The charges of the common ions in an ionic compound provide an easy way to recognize the loss and gain of electrons by the reactants, as shown in the following example (Equation 1). Iron metal and sulfur combine when heated to produce iron(II) sulfide, an ionic compound consisting of Fe²⁺ cations and S^2– anions. Iron atoms lose two electrons to form Fe²⁺ ions while sulfur atoms gain two electrons to form S^2– ions.

The loss and gain of electrons by the reactants in a chemical reaction may not be obvious from the formulas of the reactants and products if the atoms being oxidized or reduced occur in the form of a chemical compound. A method based on oxidation states has been developed to identify oxidation–reduction reactions, determine whether a substance has been oxidized or reduced, and count the electrons that are lost or gained as a result. The oxidation state may be thought of as an imaginary charge on an atom in an element or compound. Oxidation states are assigned strictly for “electron bookkeeping” purposes:

1. The oxidation state of an atom in a free element is zero.
2. The oxidation state of an atom in a monatomic ion is equal to the charge on the ion.
3. The oxidation state of fluorine in a compound is always –1.
4. The oxidation state of hydrogen in a compound is +1, except in metal hydrides (ionic compounds with metals), where it is –1.
5. The oxidation state of oxygen in a compound is –2, except in peroxides (compounds containing O—O bonds), where it is –1.
6. The sum of the oxidation states of the atoms in a neutral compound is equal to zero.
7. The sum of the oxidation states of the atoms in a polyatomic ion is equal to the charge on the ion.

A reaction is classified as a redox reaction if the oxidation states of the reactants change. Oxidation is an increase in oxidation state (corresponding to a loss of electrons). Reduction is a decrease in oxidation state (corresponding to a gain of electrons). Consider the reaction of Fe²⁺ ions with chlorine (Equation 2). The reaction is identified as a redox reaction based on the changes in oxidation states for iron and chlorine. Iron is oxidized—the oxidation state of iron increases from +2 to +3. Chlorine is reduced—the oxidation state of chlorine decreases from zero to –1. For every redox reaction, two separate half-reactions can be written. The oxidation half-reaction shows the substance that is oxidized, the product resulting from oxidation, and the number of electrons lost in the process. (The number of electrons lost is equal to the difference in oxidation states between the reactant and product.) The reduction half-reaction shows the substance that is reduced, the number of electrons gained in the process, and the product resulting from the reduction. The oxidation and reduction half-reactions for the redox reaction of Fe²⁺ with chlorine are shown below. In order to write a balanced equation for the overall redox reaction, the oxidation half-reaction must be multiplied by a factor of two. This will ensure that the number of electrons lost by Fe²⁺ is equal to or balanced by the number of electrons gained by chlorine.

Experiment Overview

The purpose of this experiment is to investigate the reactions of Fe²⁺ and Fe³⁺ ions with oxidizing and reducing agents, respectively. The results will be analyzed to determine the change in oxidation state for each reactant, the oxidation and reduction half-reactions and the balanced chemical equations for the redox reactions.

Materials

Prelab Questions

1. Identify the oxidation state of the chlorine atom in each substance shown.
   {14039_PreLab_Figure_1}
2. Potassium iodate (KIO₃) is a strong oxidizing agent and will oxidize Fe²⁺ ions to Fe³⁺. In doing so, iodate ion (IO₃^–) is reduced to elemental iodine (I₂).
   1. Use the oxidation state rules (see the Background section) to assign oxidation states to the iodine atoms in iodate ion (IO₃^–) and iodine (I₂).
   2. The half-reaction for the reduction of iodate is shown below. Use the difference in oxidation states for the iodine atoms in IO₃^– and I₂ to determine the number of electrons gained in this half-reaction. Hint: Hydrogen ions (H⁺) and water molecules (H₂O) are required to balance mass and charge.
      {14039_PreLab_Figure_2}
3. Combine the oxidation half-reaction for Fe²⁺ (see the Background section) with the reduction half-reaction for iodate (Question 2b) and write the balanced equation for the overall redox reaction of Fe²⁺ with IO₃^–. Hint: The number of electrons on each side must cancel out.

Safety Precautions

Follow all directions carefully and do not mix chemicals unless instructed in the procedure. Do not perform unauthorized reactions. Hydrochloric acid solution is a corrosive liquid. It causes severe skin burns and eye damage and is toxic by ingestion or inhalation. Sodium hypochlorite solution reacts with concentrated acids to generate poisonous chlorine gas. It is a corrosive liquid. Do not mix with acids. Potassium ferricyanide and potassium thiocyanate solutions are irritating to skin and eyes and may evolve poisonous fumes upon heating or in contact with concentrated acids. Iron(III) chloride and sodium sulfite solutions are skin and eye irritants. Avoid contact of all chemicals with eyes and skin. Wear chemical splash goggles, chemical-resistant gloves and a lab coat or chemical-resistant apron. Wash hands thoroughly with soap and water before leaving the lab.

Procedure

Reactions of Iron(II) Ions with Oxidizing Agents (Part A)

1. Place a clean, 24-well reaction plate on top of a sheet of white paper, as shown in Figure 1. Each well is identified by a unique combination of a letter and a number, where the letter refers to a horizontal row and the number to a vertical column.
   {14039_Procedure_Figure_1_Layout and numbering of a 24-well reaction plate}
2. Using a clean Beral-type pipet or eyedropper for each solution, place 20 drops of iron(II) ammonium sulfate solution into well A1 and 20 drops of iron(III) chloride solution into well A2. Record the initial color of each solution in Data Table A.
3. Add 2 drops of potassium thiocyanate solution to each well A1 and A2 and record observations.
4. Place 20 drops of iron(II) ammonium sulfate solution into each well B1, B2 and B3.
5. Add 5 drops of 3 M hydrochloric acid solution to each well B1 and B2.
6. Using a clean pipet for each solution, add the following:
   • 5 drops of hydrogen peroxide solution to well B1
   • 10 drops of potassium permanganate solution to well B2
   • 10 drops of sodium hypochlorite solution to well B3
7. Use a clean toothpick to stir each solution, if needed, and record observations.
8. Test for the presence of iron(III) ions in wells B1, B2 and B3 by adding 5 drops of potassium thiocyanate solution to each solution. Record the final color of each test mixture in Data Table A.
9. Fill in the reactants column in Data Table A.

Reactions of Iron(III) Ions with Reducing Agents (Part B)

10. Using a clean Beral-type pipet or eyedropper for each solution, place 20 drops of iron(II) ammonium sulfate solution into well C1 and 20 drops of iron(III) chloride solution into well C2. Record the initial color of each solution in Data Table B.
11. Add 2 drops of potassium ferricyanide solution to each well C1 and C2 and record observations.
12. Place 20 drops of iron(III) chloride solution into each well D1–D5.
13. Add 5 drops of 3 M hydrochloric acid and 5 drops of sodium sulfite solution to well D1. Record observations.
14. Test for the presence of iron(II) ions in well D1 by adding 2 drops of potassium ferricyanide solution. Record the final color of the solution in Data Table B.
15. Add 5 drops of sodium bromide solution to well D2. Record observations, then test for the presence of iron(II) ions by adding 2 drops of potassium ferricyanide solution. Record the final color in Data Table B.
16. Add 5 drops of sodium iodide solution to well D3. Record observations, then test for the presence of iron(II) ions by adding 2 drops of potassium ferricyanide solution. Record the final color in Data Table B.
17. Add 10 drops of Vitamin C solution to well D4. Record observations, then test for the presence of iron(II) ions by adding 2 drops of potassium ferricyanide solution. Record the final color in Data Table B.
18. Add 10 drops of pineapple juice to well D5. Record observations, then test for the presence of iron(II) ions by adding 2 drops of potassium ferricyanide solution. Record the final color in Data Table B.
19. Fill in the reactants column in Data Table B.
20. Thoroughly rinse the contents of the reaction plate down the drain with copious amounts of water. Wash the reaction plate and rinse well with distilled water.

Student Worksheet PDF

14039_Student1.pdf