JJay - 13-4-2016 at 04:57
I just got some potassium iodide from an eBay supplier that regularly screws up orders. It was really cheap, though, and the grade is supposedly
really high, so I just went ahead and bought it.
This document contains a number of purity tests for potassium iodide: http://jpdb.nihs.go.jp/jp14e/14data/Part-I/Potassium_Iodide.... But I'd like to make sure that they didn't actually send me sodium chloride or
some other substance that isn't really potassium iodide.
The best way I can think of offhand to test for iodide is by mixing a tiny sample with a little starch and water and adding a few drops of
hydrochloric acid and hydrogen peroxide and looking for a brown color... got any better ideas?Also, what is the best way to test for potassium?
Loptr - 13-4-2016 at 05:11
You will want to titrate it because there might be potassium iodide content, but it will also contain other things as well. You will want to check the
purity by determining how much of it is iodide.
I can't look it up at the moment, but there is a basic chemistry experiment that deals with iodide content in table salt. Google for that and you will
surely find a procedure. It's simple enough, and you will most likely have the needed reagents.
[Edited on 13-4-2016 by Loptr]
JJay - 13-4-2016 at 05:15
I see a procedure in that reference that requires potassium iodate, which I don't have, but I think it is available OTC....
Boffis - 13-4-2016 at 05:55
Without knowing what chemical you have available its hard to suggest a definite method, however, it you have some good quality sodium thiosulphate and
an accurate balance (at least 0.001g readability) you can perform a reasonable estimation.
The simplest method if you don't have a good standard like KIO3 is to oxidize the KI in solution and extract into an organic phase. You can use an
excess of just about any oxidizing agent that liberates iodine (H2O2, KMnO4 etc but hypochlorite may complicate the extraction) and then extract the
liberate iodine into a chlorinated solvent (if you don't have a chlorinated solvent petroleum spirit may work or other immiscible solvents). Several
extraction will be need to get practically total extraction. This is best done in a seperating funnel but you may be able to do it with a pipette.
This extract can then be titrated against a standard sodium thiosulphate solution. This should be done in a large conical flask with plenty shaking
until the organic phase is just decolourised. No indicator is required at your likely level of precision since you are really only testing to see if
its 5% or >95% KI.
JJay - 13-4-2016 at 06:10
I do have sodium thiosulfate, of course, but I'm not seeing how extracting into an organic phase would be helpful. I'm also not seeing why I would
need a milligram scale....
I'm pretty sure I can pick up uniform dosage USP potassium iodate water purification tablets at any camping store. While they are likely to contain
some impurities, I don't think the impurities would interfere with the titration.
Loptr - 13-4-2016 at 10:40
The milligram scale is able to give you a measure of very small differences, hence the ability for you to accurately discern the iodide content.
You could also just convert the all of the iodide to iodine, sublimate, and then convert it to the sodium iodide salt with sodium hydroxide. You sure
know have iodide by this point, and a pretty good estimate of purity assuming the hydroxide isn't heavily contaminated, which might be the case with
an OTC source.
[Edited on 13-4-2016 by Loptr]
aga - 13-4-2016 at 12:03
Instead of testing for what you think it actually IS, why not test for stuff that it should not contain ?
e.g. iron, sulphates etc.
Daffodile - 13-4-2016 at 12:15
If I were you, I'd just react it with an excess of Hydrochloric Acid and Hydrogen Peroxide, afterwards weighing the product. I'd also do a flame test
to see which iodide it is. Its a crude test, but it might do the trick.
An example of what I mean:
- Soak splint in salt x solution: It gives a yellow flame. Therefore your salt has the Sodium cation.
- Dissolve one and one half grams Sodium salt x in 2ml 12M Hydrochloric Acid. Add 15ml 3% Hydrogen Peroxide. Once product settles, collect it and dry
it, minimizing loss to sublimation (as Iodine is prone to doing). Divide product mass by 1.26, assuming 100% yield. Since you got 1g of product, your
Sodium Iodide is just less than 80% pure.
That would give you the purity (%) approximately, and which Alkali salt is it. If you get no product from the second step, you don't have an Iodide
salt or science hates you.
[Edited on 13-4-2016 by Daffodile]
[Edited on 13-4-2016 by Daffodile]
Boffis - 13-4-2016 at 12:38
JJ if you simply liberate the iodine with an oxidizing agent the excess oxidizing (essential unless you know how iodide is present, in which case
analysis is pointless) will simply continue to liberate iodine until it is consumed so all you will measure is the oxidiser concentration. So the way
around this is to extract the iodine (soluble in non or weakly polar solvent) from the excess oxidizing agent (if soluble in water they are probably
not very soluble in non-polar solvents). You could alway destroy the excess oxidizing agent but this is generally difficult without also removing the
iodine (you may be able to do this with H2O2 using MnO2 since it can be removed by filtration but then you would have to add something to keep the
iodine in soluble (more iodide!!!! 
or and organic solvent.
If you don't have a balance that has reasonable accuracy I wouldn't waste your with a "quantitative" estimation of iodide. Instead I would simply
carry out a qualitative test. Say add some soluble lead salt, boil until the ppt dissolves and watch the pretty gold spangles form on cooling .
I suggest you read a text such as Vogel's Textbook of Inorganic Quantitaitive Analysis; you can download it from t'internet.
@Daffodile, why are you testing for sodium the OP said its KI and it only takes a few ppm Na to give a yellow flame that masks the K flame unless you
have a set of Merwin screens.
JJay - 13-4-2016 at 14:28
It doesn't hurt to have one (unless you don't accommodate for measurement error), but you don't need a milligram balance to produce standardized
solutions unless you need five or more significant figures and are working on a microscale level.
JJay - 14-4-2016 at 12:29
I took a look at Vogel's Textbook of Quantitative Analysis, and I believe that the iodate method of titration is superior to any of the iodide
titration methods outlined in the book, assuming that standardized solutions of iodate are available or relatively easy to prepare.
It looks like it is pretty easy to obtain potassium iodate, but I'm not really sure how suitable it is as a primary standard.
Edit: It appears that potassium iodate is quite suitable as a primary standard: http://pubs.acs.org/doi/abs/10.1021/ed026p588?journalCode=jc...
[Edited on 14-4-2016 by JJay]
morsagh - 22-4-2016 at 11:30
Best will be to mix your KI (of known mass) with HCl and H2O2, than boil off H2O2 and add a little of KI to disolve iodine, than just titrate with
ascorbic acid
Sulaiman - 22-4-2016 at 13:03
I thought that KI is a very active catalyst for the decomposition of H2O2
elephant's toothpaste etc.
JJay - 22-4-2016 at 17:06
I obtained some potassium iodate. I don't have an immediate need to titrate my KI, but I will as soon as I open the package.
Oh and yeah, I do keep chloroform in stock... it is less flammable than ether and easier to make and store than DCM. I realize that not everyone has
chloroform sitting in their reagent cabinet, but you should.
[Edited on 23-4-2016 by JJay]
morsagh - 23-4-2016 at 07:10
KI can catalyze decomposition of H2O2 but only in neutral or basic solution. In acidic it is more likely to reduce H2O2 to H2O and I- to I2.
Eddygp - 25-4-2016 at 07:13
You can mix the KI with H2O2 (of known conc.) and sodium thiosulphate and titrate with KMnO4; compare with a standardised solution of KI.
EDIT: Obviously perform qualitative analysis first.
[Edited on 25-4-2016 by Eddygp]
aithecomputerguy - 26-4-2016 at 14:06
In high school, I did this with titration. If I remember correctly, take a known amount, oxidize it to iodine, then titrate with a reducing agent. I
can't remember if we used ascorbic acid or sodium thiosulfate. We also added corn starch towards the end to tell exactly when there was none left.
JJay - 26-4-2016 at 15:46
It looks to me as though the iodate titration works by oxidizing to iodine chloride.
The major advantage of iodate is that it is relatively easy to obtain in extremely pure form.