Group 7: Halogens
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What is the colour of the titration between sodium thiosulphate and potassium iodate? Is it as colourless as water?
Is interpreting the end-point a valid argument for an important error in the above reaction?
Is tranferring the potassium iodate solution using a pipette into the conical flask a major error? How about weighing 0.053 g on a 3-decimal place balance?
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Jenkin replies ...
Have you actually carried out this titration? The first essential is to understand the chemistry involved.
An excess of potassium iodide is added to the acidified solution of potassium iodate. The iodide is oxidised, liberating iodine:
IO3- + 5I- + 6H+ -> 3I2 + 3H2O
It is this iodine which is titrated with the sodium thiosulphate solution. Iodine solution is brown. As the sodium thiosulphate solution is added, the iodine is converted back to iodide ions which are colourless
I2 + S2O3 2- -> 2I- + S4O6 2-
So, during the titration the brown solution becomes paler, and at the true end-point it will just have turned colourless (like water, as you suggest). But to improve accuracy the endpoint is often ‘sharpened’ by adding a little freshly-prepared starch solution, but not until most of the iodine has reacted and the solution is a pale straw colour. On the addition of starch the solution will turn blue-black. More thiosulphate is now added, a drop at a time, until the blue-black colour has gone, so at the endpoint the solution is colourless.
Misjudging the endpoint is not likely to be a significant source of error. But the burette readings at the start and finish of the titration will be +/- 0.05 cm3, so if your titre is 20.0, it is 20.0 +/- 0.1. This is a possible error of 0.5%.
As for weighing, it will be +/- 0.001 g on a three-place balance, so 0.053 +/- 0.001 g, about 2%. If you are using the difference between two weighings it will be 4%.
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updated: 01 April 2008
