Oxidation–Reduction

Introduction

The College Board lists the following principles and concepts for oxidation–reduction in the course description for AP^® Chemistry: oxidation number, the role of the electron in redox reactions, prediction of the direction of redox reactions, and balancing these reaction equations. Use this set of integrated, interactive demonstrations to help students review the major principles of oxidation–reduction and prepare them for the AP Chemistry Exam.

Set of two demonstrations includes:

• Reactions of Iron(II) and Iron(III) Ions—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. Investigate the oxidation and reduction reactions of iron(II) and iron(III) ions, respectively.
• Colorful Oxidation States of Manganese—In nature, manganese usually occurs as an oxide. In ionic form manganese exhibits a wide range of colorful oxidation states, including +7, +6, +5, +4, +3, and +2. Display all six oxidation states in an overhead demonstration of the oxidation of manganese in its +2 form and the reduction of its +7 state.

The set of activities may be presented in a variety of ways. Each activity may be used to review a specific AP Chemistry test topic, or both demonstrations can be performed together to assess student understanding and grasp of redox concepts normally covered in the AP exam. A student worksheet is included as an optional assessment tool for the instructor.

Concepts

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

Background

Reactions of Iron(II) and Iron(III) Ions
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. 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, 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.

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 1). The reaction is identified as a redox reaction based on the change 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.

Oxidation states:

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 combining the two half-reactions to balance the equation for the overall reaction, 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.

Fe²⁺ → Fe³⁺ + e^– Oxidation half-reaction
Cl₂ + 2e^– → 2Cl^– Reduction half-reaction

Experiment Overview

Reactions of Iron(II) and Iron(III) Ions
Iron exists in the body in two forms—iron(II), Fe²⁺, and iron(III), Fe³⁺, ions. 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 oxidation and reduction of iron(II) and iron(III) compounds, respectively.

Colorful Oxidation States of Manganese
Create a series of colorful solutions of manganese-containing ions and solids to show the various hues produced by the six oxidation states of manganese.

Materials

Safety Precautions

Hydrochloric acid is a corrosive liquid and toxic by ingestion or inhalation. Sodium hypochlorite solution reacts with acids to generate 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 generate poisonous fumes upon heating or in contact with concentrated acids. Caustic and 6 M sodium hydroxide are extremely corrosive to eyes and skin. Avoid contact with eyes and skin and clean up all spills immediately. Sulfuric acid is severely corrosive to eyes, skin and other tissue 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 laboratory. Please review current Safety Data Sheets for additional safety, handling and disposal information.

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 reduced with sodium thiosulfate according to Flinn Suggested Disposal Method #12a. Save all remaining reagents in properly labeled bottles for future use.

Prelab Preparation

Reactions of Iron(II) and Iron(III) Ions

Iron(II) ammonium sulfate, Fe(NH₄)₂(SO₄)₂, 0.1 M: Use a 100-mL graduated cylinder to add about 50 mL of distilled or deionized water to the bottle of 4 g of iron(II) ammonium sulfate. Cap and mix thoroughly. Pour the solution out into a 150-mL beaker. Add the remainder of the water to the bottle, cap, and shake. Pour the water into the beaker. Prepare fresh the day of use and protect from light by wrapping the beaker with aluminum foil.

Sodium sulfite, Na₂SO₃, 0.2 M: Use a 100-mL graduated cylinder to add 100 mL of distilled or deionized water to a 150-mL beaker. Add 2.5 g of sodium sulfite to the water. Mix thoroughly. Prepare fresh the day of use.

Procedure

Reactions of Iron(II) and Iron(III) Ions

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

1. Place a clean, 6-well reaction plate on top of the overhead projector, (or on a sheet of white paper if using a flex camera), as shown in Figure 2. 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.
   {12263_Procedure_Figure_1_Layout and numbering of a 6-well reaction plate}
2. Using a clean, 10-mL graduated cylinder for each solution, add 5 mL of iron(II) ammonium sulfate solution into well A1 and 5 mL of iron(III) chloride solution into well A2. Have students record the initial color of each solution in Data Table A.
3. Using a clean, Beral-type pipet, add 5 drops of potassium thiocyanate solution to each well A1 and A2. Record observations in Data Table A.
4. Place 5 mL of iron(II) ammonium sulfate solution into each well B1, B2 and B3.
5. Using a Beral-type pipet, add 20 drops of 3 M hydrochloric acid solution to each well B1 and B2.
6. Using a clean, Beral-type pipet, add 20 drops of hydrogen peroxide solution to well B1.
7. Using a clean, 10-mL graduated cylinder for each solution, add:
   • 3 mL of potassium permanganate solution to well B2.
   • 3 mL of sodium hypochlorite solution to well B3.
8. Use a clean toothpick to stir each solution, if needed. Record observations in Data Table A.
9. 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

1. Place a clean, 6-well reaction plate on top of the overhead projector (or on a sheet of white paper if using a flex camera).
2. Using a clean, 10-mL graduated cylinder for each solution, add 5 mL of iron(II) ammonium sulfate solution into well A1 and 5 mL of iron(III) chloride solution into well A2. Have students record the initial color of each solution in Data Table B.
3. Using a clean, Beral-type pipet, add 5 drops of potassium ferricyanide solution to each well A1 and A2. Record observations in Data Table B.
4. Place 5 mL of iron(III) chloride solution into each well B1–B3.
5. Add 20 drops of 3 M hydrochloric acid and 20 drops of sodium sulfite solution to well B1. Record observations in Data Table B.
6. Test for the presence of iron(II) ions in well B1 by adding 5 drops of potassium ferricyanide solution. Record the final color of the solution in Data Table B.
7. Add 20 drops of sodium bromide solution to well B2. Record observations in Data Table B, then test for the presence of iron(II) ions by adding 5 drops of potassium ferricyanide solution. Record the final color in Data Table B.
8. Add 20 drops of sodium iodide solution to well B3. Record observations in Data Table B, then test for the presence of iron(II) ions by adding 5 drops of potassium ferricyanide solution. Record the final color in Data Table B.
9. (Optional) Clean and rinse the first 6-well reaction plate. Place 5 mL of iron(III) chloride solution into wells A1 and A2, and then add 2 mL of Vitamin C solution to well A1. Record observations in Data Table B, then test for the presence of iron(II) ions by adding 5 drops of potassium ferricyanide solution. Record the final color in Data Table B.
10. (Optional) Add 2 mL of pineapple juice to well A2. Record observations in Data Table B, then test for the presence of iron(II) ions by adding 5 drops of potassium ferricyanide solution. Record the final color in Data Table B.

Colorful Oxidation States of Manganese

Reactions of Permanganate with Reducing Agents

1. Place a clean, 6-well reaction plate on top of the overhead projector (or on a sheet of white paper if using a flex camera), as shown in Figure 2.
   {12263_Procedure_Figure_2}
2. Have students record the initial and final color and appearance of each well as reactants are added in steps 3–10. 
3. Using a clean, 10-mL graduated cylinder, add 5 mL of manganese(II) sulfate solution into wells A1, A2 and A3.
4. Using a clean 10-mL graduated cylinder, add 5 mL of potassium permanganate solution into wells B1, B2 and B3.
5. Add 5 mL of sulfuric acid solution to well A2.
6. Add 5 mL of 6 M sodium hydroxide solution to well B2.
7. Using a Beral-type pipet, add 2–3 drops of potassium permanganate solution into well A2 and stir with a toothpick.
8. Using the Beral-type pipet from step 7, add 2–3 drops of potassium permanganate solution into well A3. Do not stir.
9. Using a Beral-type pipet, add about 2 mL of sodium sulfite solution into well B2 and stir.
10. Using a clean, 10-mL graduated cylinder, add 5 mL of 50% sodium hydroxide (caustic) into well B3. Stir rapidly with a toothpick until the solution turns color.

Lab Hints

Reactions of Iron(II) and Iron(III) Ions

• 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 activity.
• 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 activity 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 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 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

• In Reactions of Iron(II) and Iron(III) Ions, a number line may be used to help students count the number of electrons corresponding to a change in oxidation state. Iodate ion, IO₃^–, for example, is a strong oxidizing agent and will oxidize Fe²⁺ to Fe³⁺. The iodate ion is reduced to elemental iodine, I₂. The oxidation state of I changes from +5 in IO₃^– to θ in I₂.
  {12263_Tips_Figure_3}
• 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.
• 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 catalyze 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.
• In Colorful Oxidation States of Manganese, while manganese has six oxidation states, it does not exist as an ion having a +5, +6 and +7 charge in its three polyatomic ion forms of MnO₄^3–, MnO₄^2– and MnO₄^–, respectively. It is the manganese in these structures that gains or loses electrons, not the entire polyatomic ion.
• The +2 oxidation state is the most stable form for manganese. Common manganese(II) compounds include manganese sulfate and manganese chloride. In aqueous solution, the manganous(II) ion exists as the Mn(H₂O)₆²⁺ ion and has an octahedral geometry. Such compounds are usually pale pink in color. The paleness of the color is a consequence of the d⁵ electron configuration. Each orbital has a single electron and any electron transitions are spin-forbidden. In nonaqueous solvents the manganous(II) ion forms numerous complexes that have a tetrahedral geometry. These tend to be much more intensely colored than the pale-pink octahedral ions. They are yellow-green in color, and some exhibit fluorescence.
• Balancing redox equations in acidic or basic solutions can be challenging for students. The following set of steps is helpful.
  • First, balance all atoms other than oxygen and hydrogen in half-cell reaction.
  • Next, balance oxygen atoms by adding water molecules to the side deficient in oxygen.
  • Now add hydrogen ions to the side deficient in hydrogen to balance these atoms. Example:
    {12263_Tips_Equation_2}
  • If the reaction occurs in basic solution, complete the first three steps as if the reaction takes place in acid. Example:
    {12263_Tips_Equation_3}
    Then add the same number of hydroxide ions as these are hydrogen ions to both sides.

    2OH^–(aq) + 2H⁺(aq) + 2e^– + ClO^–(aq) → Cl^–(aq) + H₂O(l) + 2OH^–(aq)

    Since 2OH–(aq) + 2H+(aq) → 2H2O(l), the equation becomes

    2H₂O(l) + 2e^– + ClO^–(aq) → Cl–(aq) + H₂O(l) + 2OH^–(aq) or
    H₂O(l) + 2e^– + ClO^–(aq) → Cl^–(aq) + 2OH^–(aq)

Sample Data

Reactions of Iron(II) and Iron(III) Ions

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

Part B. Reactions of Iron(III) Ions with Reducing Agents Colorful Oxidation States of Manganese Data Table

Answers to Questions

Reactions of Iron(II) and Iron(III) Ions

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).
   {12263_Answers_Table_5}
3. Fill in the blanks to show the number of electrons 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⁺ to Fe³⁺ with the appropriate half-reaction from Question 3 and write the bal-anced 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. 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.

6. 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.

      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)

7. 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.

Colorful Oxidation States of Manganese
Based on your observations and the table of manganese oxidation states, balance the following reactions.

1. In well A2, potassium permanganate is added to an acidified solution of manganese(II) sulfate. Explain the color of the product. Balance the following half-cell reaction and write the overall balanced equation for the reaction.
   {12263_Answers_Equation_4}
2. In well A3, drops of potassium permanganate are added to an acidified solution of manganese(II) sulfate. Explain the color of the product. Balance the following half-cell reaction and write the overall balanced equation for the reaction.
   {12263_Answers_Equation_5}
3. In well B2, a solution of sodium sulfite is added to a caustic potassium permanganate solution. Explain the color of the product. Balance the following half-cell reaction and write the overall balanced equation for the reaction.
   {12263_Answers_Equation_6}
4. In well B3, a highly caustic sodium hydroxide solution is added to a solution of potassium permanganate. Explain the color of the product. Balance the following half-cell reaction and write the overall balanced equation for the reaction.
   {12263_Answers_Equation_7}

Discussion

Colorful Oxidation States of Manganese
In nature, manganese usually occurs as an oxide. The primary industrial use of manganese is in the manufacture of steel. The addition of manganese to steel increases its toughness and durability. Manganese is also used to make alloys, such as manganese bronze, which is an alloy of copper, zinc and manganese. In elemental form, manganese is fairly reactive and will displace hydrogen from acids. Manganese exists in a wide range of oxidation states, including +7, +6, +5, +4, +3 and +2.

The highest oxidation state of manganese is the +7 state, correspondong to the complete removal of all the electrons from the 4s and 3d orbitals. A very good example of a compound in this oxidation state is the permanganate ion, MnO₄^–. As a solid, crystals of potassium permanganate are so intensely colored they appear black. Solutions containing the permanganate ion are also intensely colored (purple). The permanganate ion is a strong oxidizing agent and is commonly used in the chemistry laboratory for this purpose. A common oxidation–reduction titration involves the addition of a potassium permanganate solution to a solution containing the oxalate ion or the iron(II) ion. Manganese also exists in the +6 state in the form of the manganate ion, MnO₄^2–. This ion is green in color and can be obtained by reduction of the permanganate ion. It is only stable in basic solution. The +4 oxidation state of manganese is found in manganese dioxide, MnO₂, a stable dark brown solid.