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  Redox Titrations, Chemistry tutorial

Introduction:

The reactions in which substances experience changes in oxidation number are termed to as the oxidation-reduction reactions or redox reactions. Oxidation is stated as an algebraic increase in the oxidation number, or a procedure in which the electrons are lost. Reduction is stated as an algebraic reduction in oxidation number or a method in which electrons are gained. Oxidation-reduction methods should take place concurrently. The species which gains electrons is termed as the oxidizing agent, thus it is reduced. The species which loses electrons is termed as the reducing agent, thus, it is oxidized.

Theory:

Potassium permanganate, KMnO4, is a strong oxidizing agent. Permanganate, MnO4-, is an intense dark purple color. Reduction of the purple permanganate ion to the colorless Mn2+ ion, the solution will turn from dark purple to the faint pink color at equivalence point. No additional indicator is required for this titration. The reduction of permanganate needs strong acidic conditions. In this experiment, permanganate will be decreased by oxalate, C2O42- in acidic conditions. Oxalate reacts very slowly at room temperature therefore the solutions are titrated hot to make the procedure practical. The unbalance redox reaction is illustrated below.

MnO4- + C2O42- → Mn2+ + CO2 (acidic solution)

In part-I of this experiment, a potassium permanganate solution will be standardized against the sample of potassium oxalate. Once the correct normality (eq/L) of the permanganate solution is determined, it can be employed as a standard oxidizing solution. In part-II of this experiment, the standard permanganate solution will be employed to determine the concentration of iron (II) in the ferrous solution (g/L)

The unbalanced redox reaction is illustrated below.

MnO4- + Fe2+ → Mn2+ + Fe3+ (acidic solution)

Phosphoric acid will be employed to make sure that the ferric product, Fe3+ remains in its colorless form. 

Experiment:

Equipment and Reagents (Day 1) 

  • KMnO4 solid                             
  • Weighing paper                    
  • Burette 
  • 500 mL Florence Flask           
  • K2C2O4. H2O                             
  • Ring Stand 
  • Rubber Stopper                     
  • Analytical Balance           
  • Burette Clamp 
  • Hot plate or Bunsen burner      
  • 250 mL Erlenmeyer flask  
  • 6NH2SO4

Procedure (Day 1):

Part I: Preparation of a 0.1 N KMnO4 Solution. 

1) On a centigram balance, weigh around 1.0 g KMnO4 crystals on a piece of weighing paper. Add the crystals to a 500 ml Florence Flask. 

2) Add around 350 ml of distilled water to the flask. 

3) Heat the solution by occasional swirling to dissolve the KMnO4 crystals. Don't boil the solution. This might take around 30 minutes. 

4) Let the solution to cool and stopper. We will require this solution for both day 1 and day 2.

Part-II: Standardization of a KMnO4 solution:

1) On weighing paper, weigh around 0.2 to 0.3g of K2C2O4 H2O on the analytical balance. Record the correct mass. Transfer the sample to a 250 ml Erlenmeyer flask. 

2) Rinse and fill the burette by the KMnO4 solution. 

3) Add 50 ml of distilled water and 20 ml of 6 N H2SO4 to the oxalate sample in the Erlenmeyer flask. Swirl to dissolve the solids. 

4) Heat the acidified oxalate solution to around 85oC. Don't boil the solution. 

5) Record the initial burette reading. As KMnO4 solution is strongly colored, the top of meniscus might be read rather than the bottom. 

6) Titrate the hot oxalate solution by the KMnO4 solution till the appearance of the faint pink color. 

7) Record the final burette reading and compute the volume of KMnO4 utilized in the titration. 

8) Remove the titration mixture down the drain and repeat the titration by a new sample of oxalate for a net of 2 trials. 

9) The oxalic acid solution might be employed to wash the burette and the titration flask if a brown stain remains in glass-ware.

Computations:

1) By using the half-reaction method, write down a balanced redox equation for the reaction of permanganate by oxalate in the acidic solution.

2) Compute the equivalent weight of the oxalate reducing agent from the molar mass of the oxalate sample and the equivalence of electrons lost via the reducing agent in the oxidation half-reaction. 

Equivalent weight = (184 g/mol)/(# of electrons eq/mol)

3) Make use of the sample mass and the equivalent weight to compute the number of equivalents of oxalate in each and every sample.

Equivalence of reducing agent = sample mass g x eq g

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

eqred = eqox

4) Compute the normality of the KMnO4 solution from the equivalence of the oxidizing agent and the volume employed in the titration. 

5) Compute the average normality of the permanganate solution.

Equipment and Reagents (Day 2):

  • Unknown Fe2+ solution   
  • KMnO4 solution       
  • Burette Clamp 
  • 250 mL Erlenmeyer Flask   
  • 25 mL pipette     
  • Ring Stand 
  • 6 NH3PO4             
  • Pipet bulb 

Procedure (Day 2) 

Part III: Determination of the Mass of Iron in a Ferrous Solution

1) Pipet a 25 ml sample of the unknown Fe2+ solution to a 250 ml Erlenmeyer flask. 

2) Add 50 ml of distilled water and 12 ml of 6 NH3PO4 into the flask. 

3) Fill a burette by the standard KMnO4 solution and record the initial burette reading. 

4) Titrate the sample by the standard KMnO4 to a faint pink end-point and record the final burette reading. Compute the volume of KMnO4 used. 

5) Remove the ferric solution down the drain and repeat the titration by a new sample of the ferric solution for a net of 2 trials. 

6) After all trials, reject the purple permanganate solution in the suitable waste container in the fume hood. 

7) Oxalic acid might be employed to remove any brown stains left on the glass-ware.

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