Materials Included In Kit

Additional Materials Required

Prelab Preparation

The following solutions are light and air sensitive and should be prepared fresh within 1–2 days of use.

Iron(II) Ammonium Sulfate, 0.1 M: 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. Prepare fresh within days of use. This reagent is also called ferrous ammonium sulfate.

Vitamin C, 0.2 %: Dissolve a 500-mg tablet in 250 mL of water. Filter the solution to remove any insoluble residue (binder material).

Safety Precautions

Hydrochloric acid is a corrosive liquid and toxic by ingestion or inhalation. Sodium hypochlorite solution is a corrosive liquid and is moderately toxic by ingestion and inhalation. The solution reacts with acids to evolve poisonous chlorine gas. Hydrogen peroxide is a strong oxidizing agent and a skin and eye irritant. Potassium ferricyanide and potassium thiocyanate solutions are toxic by ingestion and may evolve poisonous fumes upon heating or in contact with concentrated acids. Iron(III) chloride, potassium permanganate 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 chemical-resistant apron. 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 any 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 plate 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 disposed of by reduction 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 in a typical 50-minute lab period. Review the Prelab Questions prior to lab to help students prepare for 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 the day of the lab. The solution should be pale green.
• 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).
• Other oxidizing and reducing agents that may be tested in this experiment include 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 half-reaction (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 spontaneously 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).

Teacher Tips

• 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 of 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.
• 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

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 atom in IO₃^– and I₂.

   For iodate: According to rule 5, the oxidation state of each oxygen atom is –2.
   According to rule 7: (3 –2) + (Ox. 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 below. Use the difference in oxidation states for the iodine atom in IO₃^– and I₂ to determine the number of electrons (n) gained in this half-reaction. Hint: Hydrogen ions (H⁺) and water molecules (H₂O) are required to balance mass and charge.

   IO₃^– + 6H⁺ + 5 e^– ½I₂ + 3H₂O
   The difference in oxidation state for the iodine atom in iodate ion and iodine is (+5 – 0) = 5. Each iodate ion gains five electrons (n = 5) as the oxidation state of the iodine atom is reduced from +5 to zero. 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.

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

   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.
   5Fe²⁺ 5Fe³⁺ + 5e^–
   IO₃^– + 6H⁺ + 5e^– ½I₂ + 3H₂O
   Overall reaction: 5Fe²⁺ + IO₃^– + 6H⁺ 5Fe³⁺ + ½I₂ + 3H₂O

Sample Data

Data Table A. Reaction of Iron (II) Ions with Ozidizing Agents

Data 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⁺?

   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).
   {12866_Answers_Table_3}
3. Determine the number of electrons (n) involved in each half-reaction.
   1. MnO₄^–(aq) + 8H⁺(aq) + 5e^–      Mn²⁺(aq) + 4H₂O(l)
   2. H₂O₂(aq) + 2H⁺(aq) + 2e^–         2H₂O(l)
   3. OCl^–(aq) + H₂O(l) + 2e^–           Cl^–(aq) + 2OH^–(aq)
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⁺?

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

   2. Determine the number of electrons (n) in the following half-reaction.

      SO₃^2–(aq) + H₂O(l) SO₄^2–(aq) + 2H⁺(aq) + n e^–
      The difference in oxidation states for the sulfur atom is +2. Each sulfite ion loses two electrons when it is oxidized to sulfate.
      SO₃^2–(aq) + H₂O(l) SO₄^2–(aq) + 2H⁺(aq) + 2e^–

   3. Write the balanced equation for the overall redox reaction of Fe³⁺ with a 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.

   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. Oxidation of biological molecules may be caused by several factors. (1) Oxygen and ozone in the atmosphere make the air we breathe a strongly oxidizing environment. (2) Oxidizing agents such as hydrogen peroxide and nitric oxide are normal by-products of cellular metabolism and cell processes. Vitamin C reduces and scavenges both internally and externally produced oxidants before they can react and cause problems.

11. (Optional) 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.

References

This activity was adapted from Flinn ChemTopic^™ Labs, Vol. 16, Oxidation and Reduction, Cesa. I., Editor; Flinn Scientific: Batavia, IL (2004).

Introduction

Iron exists in the body in two forms—iron(II), Fe²⁺, and iron(III), Fe³⁺ ions. Both forms of iron are important in the absorption, storage and utilization of iron by the body. Iron(II) compounds, for example, are more easily absorbed by the body, but iron is stored in the body in the form of iron(III) compounds. Iron is also an essential cofactor for many enzymes. The iron atoms in redox enzymes reversibly alternate between the +2 and +3 forms. Let’s investigate the oxidation and reduction reactions of iron(II) and iron(III) ions, respectively.

Concepts

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

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. Reduction occurs when a substance gains electrons. Because any electrons lost by one reactant must be transferred to another reactant, oxidation and reduction always occur together. 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 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.

The loss and gain of electrons by the reactants in a chemical reaction is not always obvious from the formulas of the reactants and products. A method based on oxidation states has been developed to identify oxidation–reduction reactions, to determine whether a substance has been oxidized or reduced, and to 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.

The following rules are used to assign oxidation states:

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 all the atoms in a neutral compound is equal to zero.
7. The sum of the oxidation states of all 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 (equivalent to a loss of electrons). Reduction is a decrease in oxidation state (equivalent to a gain of electrons). Consider the reaction of Fe²⁺ ions with chlorine (Equation 1). 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. The oxidation half-reaction must be multiplied by a factor of two so that the number of electrons lost by Fe²⁺ will be equal to 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 equation for the overall redox reactions.

Materials

Prelab Questions

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

   IO₃^– + 6H⁺ + n e^– ½I₂ + 3H₂O

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

Safety Precautions

Follow all directions carefully and do not perform any unauthorized reactions. Hydrochloric acid is a corrosive liquid and is toxic by ingestion or inhalation. Sodium hypochlorite solution reacts with acids to evolve poisonous chlorine gas. It is a corrosive liquid and is moderately toxic by ingestion and inhalation. Hydrogen peroxide is a strong oxidizing agent and a skin and eye irritant. Potassium ferricyanide and potassium thiocyanate solutions are toxic by ingestion and may evolve poisonous fumes upon heating or in contact with concentrated acids. 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

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

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.
   {12866_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. Record observations in Data Table A.
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
   • 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. Record observations in Data Table A.
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.

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

9. 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.
10. Add 2 drops of potassium ferricyanide solution to each well C1 and C2. Record observations in Data Table B.
11. Place 20 drops of iron(III) chloride solution into each well D1–D5.
12. Add 5 drops of 3 M hydrochloric acid and 5 drops of sodium sulfite solution to well D1. Record observations in Data Table B.
13. 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.
14. Add 5 drops of sodium bromide solution to well D2. Record observations in Data Table B, 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.
15. Add 5 drops of sodium iodide solution to well D3. Record observations in Data Table B, 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 10 drops of Vitamin C solution to well D4. Record observations in Data Table B, 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. (Optional) Add 10 drops of pineapple juice to well D5. Record observations in Data Table B, 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. Rinse the contents of the reaction plate down the drain with plenty of excess water. Wash the reaction plate and rinse well with distilled water.

Student Worksheet PDF

12866_Student1.pdf