Colorful Oxidation States of Manganese

Introduction

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

Concepts

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

Materials

Safety Precautions

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. 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. Sulfuric acid solution may be neutralized according to Flinn Suggested Disposal Method #24b. Sodium hydroxide solution may be disposed according to Flinn Suggested Disposal Method #10. Sodium sulfite solutions may be oxidized according to Flinn Suggested Disposal Method #12b. Potassium permanganate solution may be reduced according to Flinn Suggested Disposal Method #12a.

Procedure

Reactions of Permanganate with Oxidizing and 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 1.

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 each well A1, A2 and A3.
4. Using a clean 10-mL graduated cylinder, add 5 mL of potassium permanganate solution into each well B1, B2 and B3.
5. Add 5 mL of sulfuric acid solution into well A2.
6. Place 5 mL of 6 M sodium hydroxide solution into 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 a 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.

Student Worksheet PDF

12356_Student1.pdf

Teacher Tips

• This kit contains enough chemicals to perform the demonstration as written seven times: 125 mL of 0.006 M potassium permanganate solution, 50 mL of 6 M sodium hydroxide solution, 35 mL of 6 M sulfuric acid solution, 50 mL of 50% sodium hydroxide solution (caustic), 30 mL of 0.2 M sodium sulfite solution, two 6-well reaction plates, 21 Beral-type pipets, and a box of toothpicks.

• While manganese has six oxidation states, it does not exist as an ion having a +5, +6 or +7 charge. These three oxidation states of manganese are found in the polyatomic ion MnO₄^3–, MnO₄^2– and MnO₄^–, respectively. The manganese atoms in these structures lose or gain electrons when the ions are oxidized or reduced.
• The +2 oxidation state is in the most stable form for manganese. Common manganese(II) compounds include manganese sulfate and manganese chloride. In aqueous solution, the manganese(II) ion exists as the Mn(H₂O)₆²⁺ complex 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 manganese(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 hydrogens to balance these atoms.
  
  Example:
  {12356_Tips_Table_1}

  —If the reaction occurs in basic solution, complete the first three steps as if the reaction takes place in acid.
  
  Example:

  {12356_Tips_Table_2}

  Then add the same number of hydroxide ions as there 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

Answers to Questions

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 reactions and write the overall balanced equation for the reaction.

2. In well A3, potassium permanganate is added to manganese(II) sulfate. Explain the color of the product. Balance the following half-cell reactions and write the overall balanced equation for the reaction.

3. In well B2, sodium sulfite is added to a basic potassium permanganate solution. Explain the color of the product. Balance the following half-cell reactions and write the overall balanced equation for the reaction.

4. In well B3, a concentrated sodium hydroxide solution (caustic) is added to potassium permanganate. Explain the color of the product. Balance the following half-cell reactions and write the overall balanced equation for the reaction.

Discussion

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, corresponding 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 permangante 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 or black solid.