Materials Included In Kit

Additional Materials Required

Prelab Preparation

Fe²⁺ standard solution, 0.100 M, 1 L (Make fresh on day of laboratory.): Weigh out 39.210 g of Fe(NH₄)₂(SO₄)₂•6H₂O on an analytical balance and transfer to a clean 1-L volumetric flask. Fill the flask with approximately 700 mL of deionized water. Cap and mix until the solid is completely dissolved. Fill to mark with deionized water. Cap and mix thoroughly.

Oxalic acid solution, approximately 0.025 M, 1 L: Using a clean graduated cylinder, transfer 100 mL of 0.25 M oxalic acid solution to a clean 1-L volumetric flask. Fill the flask to mark with deionized water. Cap and mix thoroughly.

Potassium permanganate solution, 0.020 M, 1 L: Use a clean graduated cylinder to transfer 200 mL of 0.10 M potassium permanganate solution to a clean 1-L volumetric flask. Fill to mark with deionized water. Cap and mix thoroughly. Approximately 1 liter of 0.020 M potassium permanganate solution is needed for 12 groups of students.

Safety Precautions

Sulfuric acid (6 M) is corrosive to eyes, skin and other tissue; always add acid to water, never the reverse. Do not heat sulfuric acid. Potassium permanganate solution may be a skin irritant. The oxalic acid solution is a skin and eye irritant; it is moderately toxic by ingestion. The manganese sulfate solution is a body tissue irritant. Wear chemical splash goggles and chemical-resistant gloves and apron. Have students wash their hands thoroughly with soap and water before leaving the laboratory. Please consult current Safety Data Sheets for additional safety 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. The potassium permanganate solution may be disposed of according to Flinn Suggested Disposal Method #12a. The Part 1 and 2 titration solutions and the manganous sulfate solution may be disposed of according to Flinn Suggested Disposal Method #27f. The 6 M sulfuric acid solution may be disposed of according to Flinn Suggested Disposal Method #24b. The oxalic acid and its solution may be disposed of according to Flinn Suggested Disposal Method #24a. The ferrous ammonium sulfate may be disposed of according to Flinn Suggested Disposal Method #26a; its solution by #26b.

Lab Hints

• The titration lab teaches students how to use volumetric glassware and encourages them to develop proper laboratory techniques. Review and demonstrate these titration techniques.
• Students can expect very precise results (to 3 significant figures) that are relatively easy to obtain.
• If volumetric pipets are not available, a separate buret can be used. Be sure the buret is cleaned between uses.
• Have students approximate the amount of MnO₄^– solution needed to titrate the iron solution in Part 1 before they titrate. The approximate molarity for the MnO₄^– solution can be given as between 0.015 M and 0.025 M.
• Quantitative analysis represents a nearly invisible application of chemistry in daily life. To illustrate this, ask students how they would react if they could not trust the water they drink or medicines they take had been tested to assure their quality and safety.

Answers to Prelab Questions

1. Write the balanced net ionic equation for the reaction between MnO₄^– ions and H₂C₂O₄ in acid solution.

   First, write out each half-reaction:
   8H⁺(aq) + MnO₄^–(aq) + 5e^– → Mn²⁺(aq) + 4H₂O(l)
   2H₂O(l) + H₂C₂O₄(aq) → 2H₂CO₃(aq) + 2H⁺(aq) + 2e^–
   
   Next, balance the electrons, then add the half-reactions.
   16H⁺(aq) + 2MnO₄^–(aq) + 10e^– → 2Mn²⁺(aq) + 8H₂O(l)
   10H₂O(l) + 5H₂C₂O₄(aq) → 10H₂CO₃(aq) + 10H⁺(aq) + 10e^–
   6H⁺(aq) + 2H₂O(l) + 2MnO₄^–(aq) + 5H₂C₂O₄(aq) → 2Mn²⁺(aq) + 10H₂CO₃(aq)

2. How many moles of Fe²⁺ ions can be oxidized by 0.043 moles of MnO₄^– ions?

   From the balanced oxidation–reduction reaction, five moles of Fe²⁺ ions are oxidized by one mole of MnO₄^– ions. Therefore, the moles of Fe²⁺ ions oxidized are:

   {13824_PreLabAnswers_Equation_4}
3. 1.630 g of iron ore is dissolved in an acidic solution. This solution is titrated to a pink endpoint with 27.15 mL of a 0.020 M KMnO₄ solution.
   1. How many moles of MnO₄^– ions were consumed?
      {13824_PreLabAnswers_Equation_5}
   2. How many moles of Fe²⁺ were in the iron ore sample?

      Moles Fe²⁺ = 5 x (moles MnO₄^–)

      = 5 x 0.00054 moles
      = 0.0027 moles

   3. What is the percent of iron in the iron ore sample?

      % iron = g Fe/g sample x 100
      g Fe = moles Fe²⁺ x molar mass Fe = 0.0027 moles Fe²⁺ x 55.847 g Fe/mole = 0.15 g Fe
      % Fe = (0.15 g Fe/1.630 g sample) x 100 = 9.3%

Sample Data

Part 1. Standardization of Potassium Permanganate Solution
Molarity of Fe²⁺ ___0.100___ M

Part 2. Determination of Concentration of Oxalic Acid Solution
Molarity of MnO₄^– solution ___0.0206___ M Molarity of H₂C₂O₄ solution ___0.0260___ M

Answers to Questions

Post-Laboratory Review

1. From the Part 1 standardization data, calculate the molarity of the MnO₄^– solution for each trial. Average the values and enter the average in the Part 2 Data Table.

   Molarity of MnO₄^– = (1/5) Moles Fe²⁺/L of MnO₄^– solution

   {13824_Answers_Equation_6}
   Moles Fe²⁺(aq) = 0.00100 moles
   1. Molarity MnO₄^–(aq) = (0.200 x 0.00100 moles)/0.00965 L = 0.0207 M
   2. Molarity MnO₄^–(aq) = (0.200 x 0.00100 moles)/0.00980 L = 0.0204 M
   3. Molarity MnO₄^–(aq) = (0.200 x 0.00100 moles)/0.00960 L = 0.0208 M
   4. Average molarity = (0.021 + 0.020 + 0.021) M/3 = 0.0206 M
2. From the Part 2 titration data, calculate the molarity of the H₂C₂O₄ solution for each trial. Average the values and enter the average in the Part 2 Data Table. The balanced reaction has 5 moles of H₂C₂O₄(aq) reacting with 2 moles of MnO₄^–(aq). At the equivalence point:
   {13824_Answers_Equation_7}
   The molarity of the oxalic acid solution is equal to:
   {13824_Answers_Equation_8}
   Average molarity H₂C₂O₄(aq) = (0.0261 + 0.0258)/2 = 0.0260 M
3. How many moles of Fe²⁺ ions and MnO₄^– ions were titrated in each Part 2 trial?
   {13824_Answers_Equation_9}
4. How many moles of oxalic acid, H₂C₂O₄ were titrated in each Part 2 trial?
   {13824_Answers_Equation_10}

Introduction

A common task in analytical chemistry is the determination of the amount of a substance present in a sample or product. If the product contains a substance that can be oxidized, then it is possible to determine the number of moles of that substance by titrating the sample with a solution of a strong oxidizing agent. In this lab, an oxidizing solution will be standardized and then use to determine the number of moles of oxalic acid, a reducing agent.

Concepts

• Oxidation–reduction reaction
• Half-reaction
• Titration
• Equivalence point

Background

Oxidation–reduction reactions occur by electron transfer. The balanced chemical reaction can be written as the combination of two half-reactions, representing the oxidation reaction and the reduction reaction, respectively. For example: If solid iron is placed in a solution of gold(III) ions, the gold(III) ions are reduced to solid gold and the iron oxidized to iron(III) ions, according to the following half-reactions:

Fe(s) → Fe³⁺(aq) + 3e^– oxidation half-reaction
Au³⁺(aq) + 3e^– → Au(s) reduction half-reaction

In this experiment, potassium permanganate, KMnO₄, is used as the oxidizing agent. In an acidic solution, the MnO₄^– ion is reduced from Mn(VII) to Mn(II) according to the following half-reaction: In Part 1, a solution of KMnO₄ is standardized by titration with a solution containing a known concentration of iron(II) ions, (Fe²⁺). In the corresponding oxidation half-reaction, the Fe²⁺ ion is oxidized to Fe³⁺: Combining half-reactions 2 and 3 and balancing the number of electrons transferred gives the overall reaction equation: The balanced equation shows that 5 moles of Fe²⁺ are required to react with 1 mole of MnO₄^–. For this redox titration, the equivalence point occurs when the exact number of moles of Fe²⁺ ions has been added to react completely with all the MnO₄^– ions in solution. At this point: If the volume and molarity of the Fe²⁺ solution are known, then: Rearranging Equation 2 yields the equation for the concentration of the potassium permanganate solution: The indicator for this titration is the MnO₄^– ion itself. The MnO₄^– ion is purple in solution. At the endpoint of the titration, the solution changes from light pink to colorless.

In Part 2, the concentration of an oxalic acid solution is determined by titration with the permanganate solution standardized in Part 1. In this case, the endpoint occurs when the pink color of the MnO₄^– ion persists. The half-reaction for the oxidation of oxalic acid is: The oxidation state of carbon changes from (+3) in H₂C₂O₄ to (+4) in H₂CO₃.

Experiment Overview

The purpose of this lab is to standardize a solution of potassium permanganate by redox titration with a standard solution of iron(II) ions. A solution of oxalic acid is then titrated with the permanganate solution to determine the exact concentration of oxalic acid.

Materials

Prelab Questions

1. Write the balanced net ionic equation for the reaction between MnO₄^– ions and H₂C₂O₄ in acid solution.
2. How many moles of Fe²⁺ ions can be oxidized by 0.043 moles of MnO₄^– ions?
3. 1.630 g of iron ore is dissolved in an acidic solution. This solution is titrated to a pink endpoint with 27.15 mL of a 0.020 M KMnO₄ solution.
   1. How many moles of MnO₄^– ions were consumed?
   2. How many moles of Fe²⁺ were in the iron ore sample?
   3. What is the percent of iron in the iron ore sample?

Safety Precautions

Sulfuric acid (6 M) is corrosive to eyes, skin and other tissue; always add acid to water, never the reverse. Do not heat sulfuric acid. Potassium permanganate solution may be a skin irritant. The oxalic acid solution is a skin and eye irritant; it is moderately toxic by ingestion. The manganese sulfate solution is a body tissue irritant. Wear chemical splash goggles and chemical-resistant gloves and apron. Wash hands thoroughly with soap and water before leaving the laboratory.

Procedure

Part 1. Standardization of Potassium Permanganate Solution

1. Obtain approximately 80 mL of the potassium permanganate solution in a 100-mL beaker. Obtain 50 mL of the 0.100 M ferrous ammonium sulfate solution in another 100-mL beaker. Label both beakers.
2. Set up a clean, 50-mL buret in the ring stand and buret clamp.
3. Rinse the buret with approximately 10 mL of distilled or deionized water and then with two 5 mL portions of the MnO₄^– solution (potassium permanganate solution, KMnO₄).
4. Close the stopcock and fill the buret to above the zero mark with MnO₄^– solution.
5. Open the stopcock to allow any air bubbles to escape from the tip. Close the stopcock when the liquid level is between the 0- and 10-mL marks.
6. Record the precise level of the solution in the buret in the Part 1 Data Table. This is the initial volume of the MnO₄^– solution. (See Figure 1 for reading buret level.)
   {13824_Procedure_Figure_1}
7. With the volumetric pipet, transfer 10 mL of the 0.100 M Fe²⁺ solution to a clean 250-mL Erlenmeyer flask. Record this volume in the Part 1 Data Table.
8. Measure out 10 mL of the 6 M H₂SO₄ into a clean 10-mL graduated cylinder and add this to the Erlenmeyer flask. Swirl to mix.
9. Position the flask under the buret so the tip of the buret is within the flask but at least 2 cm above the liquid surface.
10. Titrate the ferrous ammonium sulfate solution with the MnO₄^– solution until the first trace of pink color persists for 30 seconds. Remember to swirl the flask and to rinse the walls of the flask with distilled water before the endpoint is reached.
11. Record final buret reading as the final volume of the MnO₄^– solution in the Part 1 Data Table.
12. Repeat the standardization titration two more times.

Part 2. Determination of Concentration of Oxalic Acid Solution

1. Obtain approximately 60 mL of the oxalic acid solution in a clean 100-mL beaker.
2. With a 25-mL volumetric pipet, transfer 25 mL samples of the oxalic acid solution to each of two clean 250-mL Erlenmeyer flasks. Record the volume in the Part 2 Data Table.
3. Add 5 drops of the 1.0 M MnSO₄ solution to each flask. (The Mn²⁺ ion acts as a catalyst for the reaction.)
4. Measure out 10 mL of 6 M H₂SO₄ into a graduated cylinder and add this amount to each of the 250-mL Erlenmeyer flasks. Add 20 mL of distilled water to each flask and swirl.
5. Warm the first flask to about 85 °C on the hot plate.
6. Immediately titrate this solution with the standardized MnO₄^– solution from Part 1. Record both the initial and final buret readings in the Part 2 Data Table.
7. Repeat steps 5 and 6 with the second flask.

Student Worksheet PDF

13824_Student1.pdf