Volumetric Analysis (3) Redox Titrations

The point of this blog is to introduce you to redox titrations, since the titration technique is not restricted to acid base reactions. 

Here we are thinking about using a titration to determine the stoichiometry of a reaction.

I always think that this common experiment found in most Chemistry courses in the UK and elsewhere is a bit pointless since you could just as easily balance the equation involved by the usual means.

The equation is:

Sodium thiosulphate  + Iodine  =   sodium iodide  +  sodium tetrathionite

2Na₂S₂O₃            +         I₂         =             2NaI         +       Na₂S₄O₆

so the problem is solved since the equation shows us that the stoichiometry of this reaction is 2 moles sodium thiosulphate combine with 1 mole iodine.  

But the titration is worth carrying out so as to discuss the method because it has a few quirks in it.

Of course, there is no titration curve.

But we do use an unusual indicator: starch solution.

Starch solution forms a deep blue complex with iodine. 

The starch, a few mls, is added towards the end point when the iodine solution is a golden yellow or straw colour. 

Adding starch too soon may form a deep blue–black precipitate with iodine which does not dissolve or breakdown once formed and so starch effectively removes iodine from the reaction mixture.

The end point of this titration occurs when the addition of one extra drop of sodium thiosulphate causes the deep blue colour to disappear. 

From this you can tell that the sodium thiosulphate solution is in the burette and the iodine solution is in the conical flask. 

Here is typical set up.

Here is a version of this experiment on YouTube

And another funnier version here

Here is typical set of results

Pipette solution	Iodine solution			0.500mol/dm3		10ml
Burette solution	Sodium thiosulphate			0.500mol/dm3
Indicator	Starch solution
Burette rdgs	Rangefinder	1	2	3	(4)
Final rdg (ml)	20.40	20.10	20.00	19.90
First rdg (ml)	0.00	0.00	0.20	0.00
Volume used (ml)	20.40	20.10	19.80	19.90
Mean titre (ml)	19.85

First, we see that the two nearest titres are used to give the mean value of 19.85ml

Second, we can calculate the number of moles of each reactant using  n = cV

Iodine:  n =  0.500  × 10/1000  =   0.00500moles

Sodium thiosulphate:  n = 0.500  ×  19.85/1000 =  0.009925 moles

So the ratio of moles thiosulphate to iodine is

0.009925  :  0.00500 

or  approx.  2 : 1

This result would confirm the stoichiometry of the equation we determined at the start of this blog.

Notes:

If the burette had not been initially rinsed with sodium thiosulphate, just with water, then the rangefinder would have been greater than expected as we can see.  This is because the water in the burette would dilute the first volume of thiosulphate used.

If the pipette is not rinsed with iodine solution then the water in the pipette dilutes the initial quantity of iodine so less thiosulphate is required to react with it.  This makes the rangefinder result less than it should be.

If in any titration the burette jet is not filled with sodium thiosulphate then the titre is larger than it should be so the moles thiosulphate is larger too. 

The conical flask can of course contain distilled water before the titration because the pipette transfers a given number of moles iodine to it regardless of the water in the flask.