CHAPTER 15 REDOX TITRATIONS

oxygen in the water. If acid isn’t added the solution will quickly start to turn yellow due to formation of MnO2. If this happens anyway, add some more sulphuric acid to reverse the process. The solution should be clear with a slight green tinge.

Chapter 15 – Volumetric analysis
(Redox titrations)

To prepare a standard solution of ammonium iron sulphate and use this solution to standardise a solution of potassium permanganate. Potassium permanganate (KMnO 4 ) is a powerful oxidizing agent and is a deep purple colour. It cannot be obtained in sufficiently pure form to weigh it out accurately and make up a standard solution directly. We must titrate an approximate solution of potassium permanganate (0.02 M) against a primary standard solution, ammonium iron sulphate (0.1
M).
Prepare the ammonium iron sulphate solution.
 The molecular formula of ammonium iron sulphate is NH 4 2 SO4 .FeSO4 .6H 2 O and its molecular mass is 392. The advantage of using a substance with such a large relative molecular mass is that a high degree of accuracy is assured when weighing it.
 Place a clean dry beaker on an electronic balance. Press the ‘zero’ button to bring the reading on the balance back to 0.00g.
Weigh out 9.8 g of ammonium iron sulphate crystals as accurately as possible, into the beaker.
 Dissolve the crystals by adding 20 ml of dilute sulphuric acid. Add some deionised water to the beaker. Mix the contents well.  Transfer the contents of the beaker to a
250 ml volumetric flask and make up the volume to the graduation mark, observing all of the usual procedures.
 The addition of sulphuric acid is vital as it prevents the Fe 2 ions from reacting with

Prepare the potassium permanganate solution.
 Weigh out an approximate amount of
KMnO 4 , in a beaker, enough to make an approximately 0.02 M solution.
Dissolve in dilute sulphuric acid, transfer and make up the solution in the usual manner with deionised water. (see before)



Titration procedure

Fill the burette with potassium permanganate solution following the usual steps.  Place 20 mls of ammonium iron sulphate in a conical flask using a pipette.
 The reaction which occurs during the titration can be summarized as follows,

MnO 4  8H   5Fe 2  Mn 2  5Fe 3  4H 2 O




If the reaction is to proceed as written, then plenty of sulphuric acid (20 mls) must be added to the conical flask to supply the H+ ions.
Sulphuric acid is used instead of any of the other common laboratory acids. It does not get involved in the oxidation or reduction except to donate H+ ions.

Determination of the end-point

The ammonium iron sulphate is a clear solution. As each drop of KMnO4 solution hits the conical flask it is reduced and decolourised. At the end point all the ammonium iron sulphate is used up. The next drop of KMnO4 solution turns the solution pink.

To determine the % of iron in iron tablets.
Dissolve the tablets
 Crush the tablets in a mortar and pestle.
 Add some conc.sulphuric acid to help it dissolve.  Transfer the mixture to a beaker.
 Slowly add the mixture to a beaker of deionised water.
 Rinse the pestle and mortar several times and add to beaker.
Transfer to volumetric flask
 Transfer the mixture from the beaker to the volumetric flask using a funnel.
 Rinse the beaker several times and add to the flask.
 Top up to 250 ml with deionised water in the usual manner.
Titration
 Fill the burette with 0.02 M KMnO4 solution.  Place 10 ml of the iron tablet solution into a conical flask, using a pipette.
 Titrate in the usual manner, swirling the flask, washing the sides regularly, adding slowly, drop by drop at the end point, read meniscus at eye level from the top(hard to see bottom.).
Determination of the end-point

The iron sulphate is a clear solution
(almost). As each drop of KMnO4 solution hits the conical flask it is reduced and decolourised. At the end point all the iron sulphate is used. The next drop of KMnO4 solution turns the solution pink.
Results
 Volume of permanganate = 8.4ml
 Concentration of permanganate= 0.02M
 Factor for permanganate = 1
 Volume of iron sulphate = 10ml
 Concentration of iron sulphate = ?
 Factor for iron sulphate = 5

Calculations

V1  M1
V  M2
 2
N1
N2
10  M1
8.4  0.02

5
1
8.4  0.02  5
M1 
10
M1  0.084 moles per litre

grams per litre = 0.084(152)g/L
12.77
= 12.77g = g in 250ml
4
= 3.19g in 8 tablets.
= 0.399g in 1 tablet
= 399mg per tablet
To standardise a solution of sodium thiosulphate using standard potassium permanganate. Sodium thiosulphate is an important reducing agent but it cannot be made up as a primary standard because it cannot be obtained in sufficiently pure form. For this reason a standard solution (accurately known) is made by titrating a solution of approximately the correct concentration with an iodine solution of accurately known concentration.
However, iodine cannot be made up accurately in a direct manner either. To get around this we react a standard solution of acidified potassium permanganate with excess potassium iodide solution. (See below for reactions).
Procedure
 Place 10ml of 0.02M KMnO4 into a conical flask.  Add 10ml of dilute sulphuric acid.
 Add 10 ml of 0.5M KI
 A deep red colour is produced due to liberation of iodine.

 The KMnO4 reacted with the KI in the presence of sulphuric acid as follows,
2MnO4  10I   16H   2Mn 2  5I 2  8H 2 O



The red colour is due to the presence of iodine. Titrate with sodium thiosulphate.
 Titrate this solution thiosulphate solution until the red colour has faded to pale yellow.  Add starch indicator.
 The solution goes black.
 Resume titration.
 The end point will occur suddenly as the iodine is finally used up.
 The black colour will suddenly disappear and go clear.
What happened?
 The iodine, produced by the initial reaction, reacts with the thiosulphate according to the following equation.
2

I 2  2S2O3
 S4O6
 2I 






As the iodine is used up the red fades to yellow. Starch is introduced and the titration is concluded when the black colour disappears.
Overall,
2KMnO4 = 5I2 = 10NaS2O3
Carry out the calculations as if 1KMnO4 directly with 5 NaS2O3.

Results
 Volume KMnO4 =10ml
 Molarity KMnO4= 0.02M
 Factor for KMnO4 = 1
 Volume of Thiosulphate = 9.95ml
 Molarity of thiosulphate = ?
 Factor for thiosulphate = 5

Calculations

V1  M1
V  M2
 2 n1 n2
9.95  M1
10  0.02

5
1
10  0.02  5
M1 
9.95
 0.1Molar
To determine the concentration of sodium hypochlorite in bleach.
Make up your bleach solution
 Put 25ml of bleach, with pipette, into
250ml volumetric flask.
 Add deionised water and make up to mark.  Stopper and invert 20 times to mix well.
 Prepare the burette
Prepare the burette
 Rinse the burette with the standard thiosulphate solution.
 Fill burette to above the zero mark with funnel.  Remove funnel, adjust to zero, ensure the tap connection is full.
Prepare the conical flask
 Place 10ml of the bleach solution into conical flask.
 Add 10ml dilute sulphuric acid to the flask.  Add 10ml 0.5M KI to the flask.
 Red iodine is released.
Titrate with sodium thiosulphate solution
 Titrate in the usual manner, swirling, rinsing regularly, add solution drop by drop.  When red fades to pale yellow add 4 drops of starch.
 Resume titrating until black colour clears. Results
 molarity thiosulphate = 0.1M
 volume thiosulphate = 11.2ml
 factor thiosulphate = 2
 molarity hypochlorite = ?
 volume hypochlorite = 10ml
 factor hypochlorite = 1

Calculations

V1  M1 V2  M 2
10  M1
11.2  0.1


=
n1 n2 1
2
11.2  0.1
 0.056moles per l
2  10
= 0.056  74.5 = 4.175g per l
 41.75g per l for the real soln.
= 4.175g per 100ml = 4.175% (w/v)
M1 