[Jasim]

I work in an analytical lab and occasionally I have to do a manual titration for determining if there are any peroxides present in a sample of THF (THF degrades into peroxides). I've asked a couple of my co-workers about the theory behind this reaction and didn't find an immediate answer. It seems like our method may be utilizing something similar to the iodine clock reaction. Could someone verify this and explain what is taking place for both the sample and the reference?


Here is what we do:

To a sample of THF two solutions are added at the same time - H2SO4 and potassium iodide. The sample is placed in the dark for 10 minutes. Ammonium molybdate is used as an indicator.

For reference - to a solution of potassium dichromate, KI and sodium bicarbonate are added. This is placed in the dark for 10 minutes. Starch is used as an indicator.

0.02N Sodium thiosulfate is used as the titer solution.

[Borek]

http://www.titrations.info/iodometric-titration

[Jasim]

Thank you Borek! Great site btw, looks like lots of good reading. 

[Jasim]

Forgive the messy post, I'm trying to think through this as I post.

I'm going over iodometric titrations and I'm still having trouble understanding this. I haven't looked at redox equations in years, since school. Could someone help me a little more in figuring this out?


These reactions are for iodometric titrations:

5I^- + IO^3- + 6H⁺ → 3I₂ + 3H₂O

2I^- ↔ I₂ + 2e-


The thiosulfate is used up when exposed to strong acid:
S₂O₃²⁻ (aq) + 2 H⁺ (aq) → SO₂ (g) + S (s) + H₂O

This is why I was thinking of the clock reaction.



When I combine the potassium iodine and sulfuric acid with the sample containing peroxide, this reaction occurs:
H₂O₂ + 2 KI + H₂SO₄ → I₂ + K₂SO₄ + 2 H₂O

I get iodine.

The iodine reacts with the thiosulfate titer solution in this reaction:
2S₂O₃^2- + I₂ → S₄O₆^2- + 2I^-


What's with placing the stuff in the dark? And where does the dichromate come in? My inorganic chemistry is severely lacking and I'm really rusty on redox.