UNIT 7 Redox Titrations – INORGANIC CHEMISTRY PRACTICALS

UNIT 7 Redox Titrations

8.0 Introduction

9.0 Objective

10.0 Main Content

3.1Theory

 3.2 Experiment

4.0 Conclusion

5.0 Summary

6.0 Tutor Marked Assignments

7.0 References/Further Reading

1.0 Introduction

Reactions in which substances undergo changes in oxidation number are referred to as oxidationreduction reactions or redox reactions. Oxidation is defined as an algebraic increase in oxidation number, or a process in which electrons are lost. Reduction is defined as an algebraic decrease in oxidation number or a process in which electrons are gained. Oxidation-reduction processes must occur simultaneously. The species that gains electrons is called the oxidizing agent, therefore it is reduced. The species that loses electrons is called the reducing agent, therefore, it is oxidized.

2.0 Objective

At the end of this unit should be able to perform an experiment in two parts

• To standardize a potassium permanganate solution will be against a sample of potassium oxalate.
• To use the standard permanganate solution to find the concentration of iron (II) in a ferrous solution

1. Main Content
2. Theory

Potassium permanganate,, is a strong oxidizing agent. Permanganate,, is an intense dark purple color. Reduction of purple permanganate ion to the colorless ion, the solution will turn from dark purple to a faint pink color at the equivalence point. No additional indicator is needed for this titration. The reduction of permanganate requires strong acidic conditions. In this experiment, permanganate will be reduced by oxalate, in acidic conditions. Oxalate reacts very slowly at room temperature so the solutions are titrated hot to make the procedure practical. The unbalance redox reaction is shown below.

In part I of this experiment, a potassium permanganate solution will be standardized against a sample of potassium oxalate. Once the exact normality of the permanganate solution is determined, it can be used as a standard oxidizing solution. In part II of this experiment, the standard permanganate solution will be used to find the concentration of iron (II) in a ferrous solution

The unbalanced redox reaction is shown below.

(acidic solution)

Phosphoric acid will be used to ensure that the ferric product, remains in its colorless form.

3.2 Experiment

Equipment and Reagents (Day 1)

• solid
• weighing paper
• burette
• 500 mL Florence Flask
• 
  
• Ring Stand
• Rubber Stopper
• Analytical Balance
• Burette Clamp
• Hot plate or Bunsen burner
• 250 mL Erlenmeyer flask
• 

Procedure (Day 1)

Part (I) – Preparation of a 0.1 N Solution.

1. On a centigram balance, weigh about 1.0 g crystals on a piece of weighing paper.

Add the crystals to a Florence Flask.

1. Add about of distilled water to the flask.
2. Heat the solution with occasional swirling to dissolve the crystals. Do not boil the solution. This may take about 30 minutes.
3. Allow the solution to cool and stopper. You will need this solution for both day 1 and day 2.

Part (II) – Standardization of a solution.

1. On weighing paper, weigh about of on the analytical balance. Record the exact mass. Transfer the sample to a Erlenmeyer flask.
2. Rinse and fill the burette with the solution.
3. Add of distilled water and of 6 N to the oxalate sample in the

Erlenmeyer flask. Swirl to dissolve the solids.

1. Heat the acidified oxalate solution to about. Do not boil the solution.
2. Record the initial burette reading. Because the solution is strongly colored, the top of the meniscus may be read instead of the bottom.
3. Titrate the hot oxalate solution with the solution until the appearance of a faint pink color.
4. Record the final burette reading and calculate the volume of used in the titration.
5. Discard the titration mixture down the drain and repeat the titration with a new sample of oxalate for a total of 2 trials.
6. An oxalic acid solution may be used to wash the burette and the titration flask if a brown stain remains in the glassware. Calculations

• Using the half-reaction method, write a balanced redox equation for the reaction of permanganate with oxalate in an acidic solution.
• Calculate the equivalent weight of the oxalate reducing agent from the molar mass of the oxalate sample and the equivalence of electrons lost by the reducing agent in the oxidation halfreaction.

  

• Use the sample mass and the equivalent weight to calculate the number of equivalents of oxalate in each sample.

  

At the equivalence point, the equivalence of the reducing agent is equal to the equivalence of the oxidizing agent.

• Calculate the normality of the solution from the equivalence of the oxidizing agent and the volume used in the titration.
• Calculate the average normality of the permanganate solution.

Equipment and Reagents (Day 2)

• Unknown solution
• solution
• Burette Clamp
• 250 mL Erlenmeyer Flask
• 25 mL pipet
• Ring Stand
• 
• Pipet bulb

Procedure (Day 2)

Part (III) – Determination of the Mass of Iron in a Ferrous Solution.

1. Pipet a sample of the unknown solution into a Erlenmeyer flask.
2. Add of distilled water and 1 of into the flask.
3. Fill a burette with the standard solution and record the initial burette reading.
4. Titrate the sample with the standard to a faint pink end-point and record the final burette reading. Calculate the volume of used.
5. Discard the ferric solution down the drain and repeat the titration with a new sample of the ferric solution for a total of 2 trials.
6. After all trials, discard the purple permanganate solution in the appropriate waste container in the fume hood.
7. Oxalic acid may be used to remove any brown stains left on the glassware.

4.0 Conclusion/Calculations

1. Using the half-reaction method, balance the redox reaction of permanganate with iron (II) in acidic media.
2. Calculate the equivalence of titrated.

mass concentrations for the ferrous unknown solution.

At the equivalence point, the equivalence of the oxidizing agent is equal to the equivalence of the reducing agent.

Determine the normality of the ferric reducing agent.

1. Calculate the molarity (mol/L) of the ferrous solution.

   

   

   (n = moles of electrons lost in the oxidation half-reaction.)

2. Calculate the mass concentration () of iron in the unknown solution by multiplying the molar mass of iron by the molarity of the ferrous solution.

1. Calculate the average mass concentrations for the ferrous unknown solution.

5.0 Summary

In this unit, you have been able to perform an experiment with two parts. In the first part, you standardized a potassium permanganate solution against a sample of potassium oxalate to determine the exact normality of the permanganate solution .In the second part of the experiment, you used the standard permanganate solution to find the concentration of iron (II) in a ferrous solution

6.0 Tutor Marked Assignments (TMAs)

1. Explain the following terms
   1. oxidizing agent
   2. reducing agent
2. A solution contains both iron (II) and iron(III) ions. A 50.0 mL sample of the solution is titrated with 35.0 mL of
   , which oxidizes to The permanganate ion is reduced to manganese (II) ion. Another sample of solution is treated with zinc metal, which reduces all the to
   . The resulting solution is again titrated with of

, this time is required. What are the concentrations of to in the solution?

Solution: 1) The stoichiometric relationship of permanganate to Fe(II):

is five to one.

1. Calculate moles of Fe(II) reacted:

( () =
,

() () =

1. Determine the TOTAL iron content:

(
,

( (

1. Determine Fe(III) in solution and its molarity:

   

   

2. Determine molarity of Fe(II):