Materials Included In Kit

Additional Materials Required

Prelab Preparation

To prepare 1 liter of 0.020 M potassium permanganate solution:

1. Using a clean graduated cylinder, transfer 200 mL of 0.10 M KMnO₄ solution to a clean 1000-mL volumetric flask.
2. Fill to mark with deionized water.
3. Cap and mix thoroughly.
4. Approximately 1 liter of 0.020 M potassium permanganate solution is needed for 12 student groups.

Safety Precautions

Sulfuric acid and sodium hydroxide are severely corrosive to eyes, skin and other body tissues. Always add acid to water, never the reverse. Keep sodium carbonate or sodium bicarbonate on hand to neutralize acid spills. Dilute potassium permanganate solution is a skin and eye irritant and a strong stain—it will stain skin and clothing. Avoid contact of all chemicals with eyes and skin. Wear chemical splash goggles, chemical-resistant gloves and a chemical-resistant apron. Remind students to wash their 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. Excess potassium permanganate solution may be reduced by adding hydrogen peroxide to the solution. This will convert the MnO₄^– ion to solid MnO₂. Separate the resulting mixture by filtration and place the solid MnO₂ in the trash according to Flinn Suggested Disposal Method #26a. The solutions remaining after the titrations are complete are acidic and contain Mn²⁺ ions. Neutralize the acidic solutions with base according to Flinn Suggested Disposal Method #24b, and add extra hydrogen peroxide (about 5 mL) to the solutions to convert the Mn2+ ions to MnO2, an insoluble brown solid. Separate the resulting mixture by filtration and place the solid MnO2 in the trash according to Flinn Suggested Disposal Method #26a. The 3 M sulfuric acid solution may be neutralized according to Flinn Suggested Disposal Method #24b. The solid ferrous ammonium sulfate may be placed in the trash according to Flinn Suggested Disposal Method #26a and its solution rinsed down the drain with excess water according to Flinn Suggested Disposal Method #26b.

Lab Hints

• Remind students to consider which reactant volumes require high precision measurements and which do not. For a titration, the volumes of reacting solutions must be known with higher precision to give more accurate results. Other species not directly involved in the analysis do not require an exact volume measurement (e.g., sulfuric acid solution, water).
• Remind students about proper quantitative transfer of the solid ferrous ammonium sulfate. Demonstrate the proper technique for rinsing the weighing dish into the flask to be sure all of the solid is transferred.
• A potassium permanganate solution is not a primary standard solution because the solid material usually contains impurities, such as nitrates, chlorides, chlorates and sulfates. In order to accurately determine the precise concentration of the solution, the potassium permanganate must be titrated with a primary standard. Two possible primary standards for potassium permanganate are ferrous ammonium sulfate [Fe(NH₄)₂(SO₄)₂] and oxalic acid (H₂C₂O₄). Although ferrous ammonium sulfate oxidizes in air, it provided reproducible and accurate standardization measurements when tested in our lab.
• Hydrogen peroxide is a stronger oxidizing agent than potassium permanganate. The standard reduction potential is 1.76 V for hydrogen peroxide versus 1.51 V for permanganate ion. In this experiment, however, hydrogen peroxide is oxidized by potassium permanganate. This is because the permanganate ion cannot be further oxidized—the Mn atom in MnO₄^– is already in the highest oxidation state possible (+7). Hydrogen peroxide can be oxidized to oxygen gas. Although hydrogen peroxide is typically regarded for its oxidizing ability, hydrogen peroxide can act as both an oxidizing agent and a reducing agent, as seen in this experiment. The well known decomposition of hydrogen peroxide to give oxygen and water is an example of a disproportionation reaction in which one hydrogen peroxide molecule is oxidized, the other is reduced.
• Have students estimate the amount of MnO₄^– solution needed to titrate the iron solution in Part 1 before they titrate. The approximate molarity of MnO₄^– can be given as between 0.015 M and 0.025 M.
• Increase the “real-world” application or value of this experiment by asking students to bring in old bottles of hydrogen peroxide from home to analyze. The typical drugstore product is labeled with an expiration date. It is not unusual for students to find bottles of hydrogen peroxide well beyond the expiration date in their medicine cabinets. These are great samples to analyze to see if the quality of the solution has deteriorated with age.

Further Extensions

Opportunities for Undergraduate Research
Hydrogen peroxide is found in many common household products, including teeth whitening strips, contact lens cleaner, hair dyes and mouth cleansers. The labels for such products may list the percent hydrogen peroxide in the product. Research a product containing hydrogen peroxide and confirm its use. Design an experiment to determine the actual percentage of hydrogen peroxide in each product.

Direct your students to the website listed below to search for products containing hydrogen peroxide. The website is hosted by the U.S. Department of Health and Human Services and lists many products with hydrogen peroxide as an ingredient. Students should pick a product readily available to them, research how the other ingredients may interact with potassium permanganate and perform the titration.

http://hpd.nlm.nih.gov/cgi-bin/household/brands?tbl=chem&id=32

Answers to Prelab Questions

1. Balance the oxidation and reduction half-reactions for the iron(II) ion and permanganate ion, respectively, and write the balanced chemical equation for the overall reaction between Fe²⁺ and MnO₄^– in acidic solution.

   5e^– + MnO₄^– + 8H⁺ → Mn²⁺ + 4H₂O
   Fe²⁺ → Fe³⁺ + e^–
   8H⁺(aq) + MnO₄^–(aq) + 5Fe²⁺(aq) → Mn²⁺(aq) + 5Fe³⁺(aq) + 4H₂O (l)

2. How many moles of Fe²⁺ ions can be oxidized by 0.067 moles of MnO₄^– ions?

   0.067 mol MnO₄^– x 5 mol Fe²⁺/1 mol MnO₄^– = 0.335 mol Fe²⁺

3. A sample of oxalic acid, H₂C₂O₄, is titrated with a standardized solution of KMnO₄. A 25.00-mL sample of oxalic acid required 12.70 mL of 0.0206 M KMnO₄ to achieve a pink colored solution. The balanced equation for this reaction is shown:

   6H⁺(aq) + 2MnO₄^–(aq) + 5H₂C₂O₄(aq) → 10CO₂(g) + 8H₂O(l) + 2Mn²⁺(aq)

   1. What does the pink color signify in this reaction?

      The pink color signifies the equivalence point and the end of the reaction.

   2. What is the mole ratio of H₂C₂O₄ molecules to MnO₄^– ions?

      The mole ratio is 5 mol H₂C₂O₄:2 mol MnO₄^–.

   3. How many moles of MnO₄^– ions reacted with the given amount of H₂C₂O₄?

      0.0206 M KMnO₄ x 0.01270 L = 2.62 x 10^–4 mol KMnO₄ = 2.62 x 10^–4 mol MnO₄^–

   4. How many moles of H₂C₂O₄ were present in the sample?

      2.62 x 10^–4 mol MnO₄^– x 5 mol H₂C₂O₄/2 mol MnO₄^– = 6.54 x 10^–4 mol H₂C₂O₄

   5. What is the molarity of the H₂C₂O₄ solution?

      M = mol/L = 6.54 x 10^–4 mol H₂C₂O₄/0.02500 L = 0.0262 M H₂C₂O₄

   6. If the density of the oxalic acid solution is approximately 1.00 g/mL, what is the percentage by mass of oxalic acid in the solution?

      Mass H₂C₂O₄ = 6.54 x 10^–4 mol x 90.03 g/1 mol = 0.0589 g H₂C₂O₄
      Mass of solution = 25.00 mL x 1.00 g/1 mL = 25.0 g
      %H₂C₂O₄ = 0.0589 g/25.0 g x 100 = 0.236% H₂C₂O₄

Sample Data

Introductory Activity

*Ferrous ammonium sulfate is provided as a hexahydrate, Fe(NH₄)₂(SO₄)₂•6H₂O, molar mass = 392.16 g/mole.

Sample Calculations
Calculated number of moles of KMnO₄ reacted in each trial in Part A.

Moles Fe²⁺ = Mass Fe(NH₄)₂(SO₄)₂ titrated/Molar mass Fe(NH₄)₂(SO₄)₂
Moles Fe²⁺ = 0.393 g/392.16 g/mol = 0.00100 mol Fe²⁺
0.00100 mol Fe²⁺ x 1 mol MnO₄^–/5 mol Fe²⁺ = 2.00 x 10^–4 mol MnO₄^–

Calculated concentration of KMnO₄ in each trial in Part A.

Molarity KMnO₄ = Moles KMnO₄/Volume KMnO₄ added
[KMnO₄] = 2.00 x 10^–4 mol/0.01062 L = 0.0188 M KMnO₄

Results for Analysis of Hydrogen Peroxide Sample Calculations
Calculated number of moles of H₂O₂ reacted in each trial in Part B.

Moles KMnO₄ = Molarity KMnO₄ x Vol. KMnO₄ added
Moles KMnO₄ = 0.0188 M x 0.01846 L = 3.47 x 10^–4 mol KMnO₄
Moles H₂O₂ = Moles KMnO₄ x (5 mol H₂O₂/2 mol KMnO₄)
Moles H₂O₂ = 3.47 x 10^–4 mol KMnO₄ x (5 mol H₂O₂/2 mol KMnO₄) = 8.68 x 10^–4 mol H₂O₂

Calculated mass of H₂O₂ reacted in each trial of Part B.

Mass H₂O₂ = Moles H₂O₂ x Molar Mass H₂O₂
Mass H₂O₂ = 8.68 x 10^–4 mol H₂O₂ x 34.01 g/mol = 0.0295 g H₂O₂
Assuming the density of the hydrogen peroxide solution was 1.00 g/mL, calculate the percent by mass of H₂O₂ in solution in each trial of Part 2.
%H₂O₂ = (Mass H₂O₂/Mass solution) x 100
%H₂O₂ = (0.0295 g H₂O₂/1.00 g solution) x 100 = 2.95%

Calculate the average percentage of hydrogen peroxide in Part B.

(2.95% + 2.97%)/2 = 2.96%
The average value for the percent hydrogen peroxide in the bottle was 2.96%. This agrees with the concentration reported on the bottle (3%). The experimental percentages may be lower than the reported percentage for several reasons. Hydrogen peroxide decomposes over time and, depending on the age of the bottle, could result in a lower experimental percentage. The lower calculated percentage may be attributed to not carrying the titration out to the equivalence point. By not reaching the equivalence point, the amount of the permanganate ion used would be less resulting in a smaller calculated amount of hydrogen peroxide.

Answers to Questions

Guided-Inquiry Design and Procedure

1. Balance the oxidation–reduction half reactions for hydrogen peroxide and permanganate ion, respectively. Write the net ionic equation for the reaction between MnO₄^– ions and H₂O₂ in acidic solution.

   8H⁺(aq) + MnO₄^–(aq) + 5e^– → Mn²⁺(aq) + 4H₂O(l)
   H₂O₂(aq) → O₂(g) + 2H⁺(aq) + 2e^–
   6H⁺(aq) + 2MnO₄ –(aq) + 5H₂O₂(aq) → 5O₂(g) + 2Mn²⁺(aq) + 8H₂O(l)

2. Asuming the density of 3% H₂O₂ is 1.00 g/mL, what mass of H₂O₂ is in a 1.00 mL sample? How many moles is this?

   1.00 mL solution x 1.00 g/mL = 1.00 g solution
   1.00 g solution x 0.03 = 0.0300 g H₂O₂
   0.0300 g H₂O₂ x 1 mol/34.0 g = 8.82 x 10^–4 mol H₂O₂

3. Based on the stoichiometry of the reaction between MnO₄^– ions and H₂O₂ in acidic solution, estimate the volume of 0.02 M MnO₄^– solution required to titrate a 1-mL sample of 3% H₂O₂.

   1.00 mL solution x 1.00 g/mL = 1.00 g solution
   1.00 g solution x 0.03 = 0.0300 g H₂O₂
   0.0300 g H₂O₂ x 1 mol/34.0 g = 8.82 x 10^–4 mol H₂O₂
   8.82 x 10^–4 mol H₂O₂ x 2 mol MnO₄^–/5 mol H₂O₂ = 3.53 x 10^–4 mol MnO₄^–
   M = mol/L
   L = moles MnO₄^–/Molarity MnO₄^–
   L = 3.53 x 10^–4 mol/0.02 M = 0.0176
   L = 17.6 mL MnO₄^– solution

4. In the Introductory Activity, the redox titration was performed in an acidic solution. Research the products of the redox reaction between MnO₄^– ions and H₂O₂ in a basic solution. Would these products affect your ability to accurately judge the endpoint of the titration?

   If the oxidation–reduction reaction of the permanganate ion is performed in a basic or neutral solution, the product will be an insoluble brown solid, manganese dioxide (MnO₂). The balanced half-reaction is: 2H₂O + MnO₄^– + 3e^– → MnO₂ + 4OH^–. The formation of a solid may hinder the ability to correctly determine when the endpoint is reached because there may be no clear color indication.

5. Is it necessary to know the exact volume of: (a) hydrogen peroxide solution added to the flask? (b) rinse water added to the flask during the filtration? (c) potassium permanganate solution added to the flask?
   1. It is necessary to know the exact volume of hydrogen peroxide because that is the species being studied in the reaction. If the precise volume of the solution is not known, it will limit the precision of the results to fewer significant figures.
   2. It is not necessary to know the exact volume of water added to the flask because it is not participating in the reaction and the amount of hydrogen peroxide is unaffected by the amount of water in the sample.
   3. It is necessary to know the exact volume of potassium permanganate because it is an active reactant in the reaction and is used to determine the amount of hydrogen peroxide in the sample through stoichiometric calculations.
6. Write a detailed step-by-step procedure for a titration experiment to determine the percent H₂O₂ in a sample: Include all the materials, glassware and equipment that will be needed, safety precautions that must be followed, the concentrations of reactants, etc. Using a 10-mL graduated cylinder, 10 mL of deionized water is added to a clean 250-mL flask.

   Using a clean 10-mL graduated cylinder, 10 mL of 3 M H₂SO₄ is added to the water in the flask. Clean a 1-mL serological pipet with small portions of the 3% H₂O₂ solution. Measure 1.0 mL of the H₂O₂ solution, record this volume on an appropriate data table, and add it to the 250-mL flask containing the water and H₂SO₄. Swirl to mix completely. Fill the buret with the MnO₄^– solution so the volume reading is between 0- and 10-mL. Record the actual volume in an appropriate data table. Position the flask under the buret so that the tip of the buret is within the flask but at least 2 cm above the liquid surface. Titrate the H₂O₂ solution with the MnO₄^– solution until the first trace of pink color persists for 30 seconds. Swirl the flask and rinse the walls of the flask with distilled water before the endpoint is reached. Record the final buret reading as the final volume of the MnO₄^– solution in an appropriate data table. Repeat the titration one more time.

Post-Laboratory Review
Examine the five reactions shown and identify those that can be classified as oxidation–reduction.

1. 2H₃PO₄ + 3Ca(OH)₂ → Ca₃(PO₄)₂ + 6H₂O

   This is not an oxidation–reduction reaction.

2. 2Cr + 3Cl₂ → 2CrCl₃

   This is an oxidation–reduction reaction.

3. C₆H₁₂O₆ + 6O₂ → 6CO₂ + 6H₂O

   This is an oxidation–reduction reaction.

4. Na₂CO₃ → Na₂O + CO₂

   This is not an oxidation–reduction reaction.

5. 2VO²⁺ + Zn + 4H+ → 2V³⁺ + Zn²⁺ + 2H₂O

   This is an oxidation–reduction reaction.

The decomposition of a compound into simpler substances by means of an electrical current is called electrolysis.

6. Write the balanced chemical equation for the electrolytic decomposition of water to its elements.

   2H₂O(l) → 2H₂(g) + O₂(g)

7. Balance the following oxidation and reduction half-reactions for the decomposition of water.

   ___H₂O → ___O2 + ___H⁺ + ___e^–
   ___H₂O + ___e^– → ___H₂ + ___OH^–

8. Explain how the oxidation and reduction half-reactions may be combined to give the balanced chemical equation for the decomposition of water. What happens to the electrons and to the H⁺ and OH^– ions?

   The electrons must balance or “cancel out” when the oxidation and reduction half-reactions are combined. The reduction half-reaction must therefore be multiplied by a factor of two. H⁺ and OH^– ions generated in the individual half-reactions combine to form water molecules.

The usefulness of metals in structural applications depends on their physical and chemical properties. Corrosion is the oxidation of metals and is a common failing point for a metal structure.

9. Examine the observations above for the metals Cu, Mg and Zn. Identify the metal that is most susceptible to oxidation (corrosion). Identify the metal that is most resistant to oxidation.

   Magnesium is the metal most susceptible to corrosion and oxidation because it reacted with the most metal ion solutions. Copper is the metal most resistant to corrosion and oxidation because it reacted with the fewest metal ion solutions.

10. Because silver metal is expensive, it was not used in the tests shown above. Based on the reactions of Cu, Mg and Zn with silver nitrate, explain why it was not necessary to test silver metal in order to deduce its reactivity.

    All of the metals tested reacted with silver nitrate. These reactions only occurred in one direction (e.g., Mg reacted with Zn(NO₃)₂, but Zn did not react with Mg(NO₃)₂). Therefore, it can be concluded that that solid silver, Ag, would not react with any of the metal ion solutions.

Introduction

Hydrogen peroxide is typically regarded as an “environmentally friendly” alternative to chlorine for water purification and wastewater treatment. Hydrogen peroxide readily decomposes in the presence of heat, light and catalysts. The quality of a hydrogen peroxide solution must be regularly checked to ensure its effectiveness. The concentration of hydrogen peroxide can be analyzed by an oxidation–reduction titration with potassium permanganate.

Concepts

• Oxidation–reduction
• Titration
• Standardization
• Half-reactions
• Equivalence point
• Percent composition

Background

Titration is a method of volumetric analysis—the use of volume measurements to analyze the concentration of an unknown. The most common types of titrations are acid–base titrations, in which an acid, for example, is analyzed by measuring the amount of standard base solution required to neutralize a known amount of the acid. A similar principle applies to oxidation– reduction reactions. If a solution contains a substance that can be oxidized, then the concentration of that substance can be analyzed by titrating it with a standard solution of a strong oxidizing agent.

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. In an overall, balanced redox reaction, the number of electrons lost by the species being oxidized is always equal to the number of electrons gained by the species being reduced.

Potassium permanganate, KMnO₄, is a common oxidizing agent used as a titrant in redox titrations. In an acidic solution, the MnO₄^– ion is reduced according to the following unbalanced half-reaction:

Potassium permanganate is not considered a primary standard for analytical purposes. Common impurities include chlorine in the form of chloride and chlorate ions, nitrogen compounds, and sulfur as sulfate. In order to accurately determine the concentration of a KMnO₄ solution, it may be titrated against a solution containing a known concentration of iron(II) ions, Fe²⁺. Ferrous ammonium sulfate, Fe(NH₄)₂(SO₄)₂, serves as a primary standard to titrate the unknown KMnO4 solution. In the corresponding half-reaction, the Fe²⁺ ion is oxidized to Fe³⁺. For this redox titration, the equivalence point occurs when the exact number of moles of MnO₄^– ions has been added to react completely with all the Fe²⁺ ions in the solution of the primary standard. The indicator for this titration is the MnO₄^– ion itself. The MnO₄^– ion is purple in solution and its reduction product, Mn²⁺, is almost colorless. At the endpoint of the titration, the solution changes from colorless to light pink as the last drop of MnO₄^– added does not react and keeps its color.

A solution of hydrogen peroxide will be titrated with the standardized potassium permanganate solution to determine the H₂O2 concentration. The endpoint occurs when the pink color of the MnO₄^– ion persists. The unbalanced half-reaction for the oxidation of hydrogen peroxide is:

Experiment Overview

The purpose of this advanced inquiry investigation is to determine the percent composition of a common “drug store” bottle of hydrogen peroxide through an oxidation–reduction titration with potassium permanganate. The lab begins with an introductory activity to standardize a solution of potassium permanganate by redox titration against a primary standard, ferrous ammonium sulfate. This standardization procedure provides a model for guided-inquiry design of an experiment to determine the percent hydrogen peroxide in a sample. Additional products containing hydrogen peroxide may be analyzed as part of optional extension activities.

Materials

Prelab Questions

1. Balance the oxidation and reduction half-reactions for the iron(II) ion and permanganate ion, respectively, and write the balanced chemical equation for the overall reaction between Fe²⁺ and MnO₄^– in acidic solution.
2. How many moles of Fe²⁺ ions can be oxidized by 0.067 moles of MnO₄^– ions? 3. A sample of oxalic acid, H₂C₂O₄, is titrated with a standardized solution of KMnO₄. A 25-00 mL sample of oxalic acid required 12.70 mL of 0.0206 M KMnO₄ to achieve a pink-colored solution. The balanced equation for this reaction is shown.

   6H⁺(aq) + 2MnO₄^–(aq) + 5H₂C₂O₄(aq) → 10CO₂(g) + 8H₂O(l) + 2Mn²⁺(aq)

   1. What does the pink color signify in this reaction?
   2. What is the mole ratio of H₂C₂O₄ molecules to MnO₄^– ions?
   3. How many moles of MnO₄^– ions reacted with the given amount of H₂C₂O₄?
   4. ow many moles of H₂C₂O₄ were present in the sample?
   5. What is the molarity of the H₂C₂O₄ solution?
   6. If the density of the oxalic acid solution is approximately 1.00 g/mL, what is the percentage by mass of oxalic acid in the solution?

Safety Precautions

Sulfuric acid is corrosive to eyes, skin and other tissue; always add acid to water, never the reverse. Notify your instructor and clean up all acid spills immediately. Dilute potassium permanganate solution is a skin and eye irritant and strong stain—it will stain skin and clothing. 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 follow all laboratory safety guidelines.

Procedure

Introductory Activity

Standardizing Potassium Permanganate Solution

1. Obtain approximately 80 mL of the potassium permanganate solution in a 100-mL beaker. Label the beaker.
2. Set up a clean 50-mL buret in a buret clamp on a support stand.
3. Rinse the buret with approximately 10 mL of distilled or deionized water and then with two 5-mL portions of the potassium permanganate solution, KMnO₄ (MnO₄^–).
4. Close the stopcock and fill the buret to just above the zero mark with the MnO₄^– solution.
5. Open the stopcock to allow any air bubbles to escape from the buret 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 an appropriate data table. This is the initial volume of the MnO₄^– solution.
7. Obtain a mass between 0.4–0.5 g of ferrous ammonium sulfate [Fe(NH₄)₂(SO₄)₂•6H₂O] in a clean weighing dish using a milligram balance. Record the precise mass in an appropriate data table.
8. Measure 10 mL of the 3 M H₂SO₄ solution into a clean 10-mL graduated cylinder. Measure 10 mL of distilled or deionized water into a separate, clean 10-mL graduated cylinder. Add these to a clean 250-mL Erlenmeyer flask. Swirl to mix.
9. Add the solid ferrous ammonium sulfate to the flask. Swirl the flask to dissolve the solid. Hint: Be sure to transfer all of the solid to the flask. Rinse the weighing dish with distilled water into the flask to capture any lingering solid.
10. Position the flask under the buret so that the tip of the buret is within the flask but at least 2 cm above the liquid surface.
11. 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 rinse the walls of the flask with distilled water before the endpoint is reached.
12. Record final buret reading as the final volume of the MnO₄^– solution in an appropriate data table.
13. Repeat the standardization titration one more time.

Analyze the Results
Use the collected volume data to determine the precise (accurate) concentration of potassium permanganate. This should be done before moving on to Guided-Inquiry Design and Procedure.

Guided-Inquiry Design and Procedure

Determination of Percent Hydrogen Peroxide
Form a working group with other students and discuss the following questions.

1. Balance the oxidation–reduction half reactions for hydrogen peroxide and permanganate ion, respectively. Write the net ionic equation for the reaction between MnO₄^– ions and H₂O₂ in acidic solution.
2. Assuming the density of 3% H₂O₂ is 1.00 g/mL, what mass of H₂O₂ is in a 1.00-mL sample? How many moles is this?
3. Based on the stoichiometry of the reaction between MnO₄^– ions and H₂O₂ in acidic solution, estimate the volume of 0.02 M MnO₄^– solution required to titrate a 1-mL sample of 3% H₂O₂.
4. In the Introductory Activity, the redox titration was performed in an acidic solution. Research the products of the redox reaction between MnO₄^– ions and H₂O₂ in a basic solution. Would these products affect your ability to accurately judge the endpoint of the titration?
5. Is it necessary to know the exact volume of: (a) hydrogen peroxide sample added to the flask? (b) rinse water added to the flask during the titration? (c) potassium permanganate solution added to the flask?
6. Write a detailed step-by-step procedure for a titration experiment to determine the percent H₂O₂ in a sample. Include all the materials, glassware and equipment that will be needed, safety precautions that must be followed, the concentrations of reactants, etc.
7. Review additional variables that may affect the reproducibility or accuracy of the experiment and how these variables can be controlled.
8. Carry out the experiment and record the results in an appropriate data table.

Analyze the Results
Use the collected volume data to determine the percentage of hydrogen peroxide in your sample. Assume the density of the hydrogen peroxide solution is 1.00 g/mL.

Student Worksheet PDF

13834_Student1.pdf