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Assured Fish
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sad.gif posted on 18-7-2017 at 18:33
What have i done


Ok so today after work i decided to prepare some elemental iodine via the classic way of mixing HCl with a potassium iodine solution and adding hydrogen peroxide.
I had done this before about a year ago, except on a much smaller scale and only using 3% H2O2, so i wasn't at all worried about something going wrong.
Unfortunately I only have concentrated hydrogen peroxide on hand which is around 28%, i didn't think anything of it but my bottle contained exactly 200mls of this 28% H2O2 and so foolishly i decided to use the entire lot, thinking that a reaction between H2O2 and iodine would not occur.

So I dissolved roughly 50g of potassium iodide in 100mls of dH2O and then added 100mls of 28% HCl, the solution turned yellow and potassium chloride precipitated as expected.
I then added this solution to 200mls of my 28%ish hydrogen peroxide and iodine immediately precipitated as expected.
I then left the solution outside wile i went to set up for vacuum filtration, when i came back though there was a lot of bubbling occurring and what looked like iodine vapor evaporating out of solution, I took a closer look and i saw the iodine was being eaten away in the solution.
The reaction got more vigorous and i had to stand back, it started producing an orange foam and producing more iodine vapor and i had to leave the immediate area to avoid the fumes.
Eventually the reaction died down and i was left with what you see below.

IMG_20170719_142357.jpg - 884kB

What have i done, did the excess peroxide violently decompose destroying my iodine in the process, did it somehow form some oxygen iodine complex which quickly evaporated out of solution.
I am at a loss. 50g of KI down the drain. Hopefully i can still scavenge a lesson from this.
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[*] posted on 18-7-2017 at 20:13


Why didn't you just dilute the H2O2 ?
I don't know what happend, but what I can tell you us that 200mls of that conc. H2O2 was WAY too much !
It might have been the heat that sumblimed the iodine ?
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[*] posted on 18-7-2017 at 20:57


Because im an idiot, i started again and used dilute H2O2 and i managed to get alright yield but its still wet at the moment.

Unfortunately I did not think about it properly before beginning, Im skeptical of it just being heat that sublimed the iodine because the solution did not even boil, I was still able to handle the beaker easily so i doubt the solution exceeded 50*C.

Wile doing it a second time, i had a similar issue with the filtrate, as i transferred it to another beaker, the solution decomposed producing a little iodine vapor, thankfully i didn't loose any iodine this time. I had diluted it down to around 10% and added no more than 100mls of this solution, equivalent to around 0.3 mols, which should be just enough to oxidize the iodide with a little chloride side oxidation.
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[*] posted on 18-7-2017 at 21:07


The classic elephant toothpaste experiment the iodide decomposed the hydrogen peroxide most of the time 3 percent solution is used also the temperature can effect the reaction and decompose the peroxide. Some hypochlorous acid is created by the reaction of hydrogen peroxide and hydrochloric acid the high temperature the iodine boiled out of solution
The color looks odd to me ah found it its Potassium Triiodide is observed to be a red colour in solution which explains why the iodide seem to disappear.

[Edited on 19-7-2017 by symboom]




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[*] posted on 19-7-2017 at 17:26


It could be that the concentrated H2O2 oxidized the iodine all the way up to iodate IO3-, which is water soluble. If that's the case, you can precipitate it with a calcium salt as calcium iodate (Ca(IO3)2). Just dump some CaCl2 solution into it and see if you get a white precipitate.
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[*] posted on 19-7-2017 at 19:50


Quote:

It could be that the concentrated H2O2 oxidized the iodine all the way up to iodate IO3-, which is water soluble. If that's the case, you can precipitate it with a calcium salt as calcium iodate (Ca(IO3)2). Just dump some CaCl2 solution into it and see if you get a white precipitate.

I followed your advice but no precipitate formed beyond the CaCl pellets that had not dissolved.

However I think i may have found the answer in another thread.
http://www.sciencemadness.org/talk/viewthread.php?tid=65425
Apparantly it is possible to get an aqueous solution of IClI (iodine monochloride), This actually makes sense as the vapor coming out of solution is an orange color.
Its a pity i don't have any ethyl acetate or ether on hand.
The good thing is that this is reversible with the addition of NaOH as highlighted by the last post by woelen in the following offshoot thread that was on the first thread i referenced.
http://www.sciencemadness.org/talk/viewthread.php?tid=24862

IClI + NaOH -----> NaOCl + I2

In the following photo you see the solution turned dark and an iodine suspension was seen after the addition of NaOH, although this is difficult to see as i took the photos with my phone.


IMG_20170720_153552[1].jpg - 669kB


IMG_20170720_153543[1].jpg - 876kB
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[*] posted on 19-7-2017 at 20:05


Unfortunately now im just left with a second problem, as the iodine precipitate has gone and reacted with the hypochlorite solution and formed a yellow solution.
I don't suppose anyone knows what has happened here, all i could find on this reaction was a bunch or religious nut jobs using it to demonstrate that drinking bleach can remove your sins.
Has the iodine formed sodium hypoiodite with the excess NaOH?

Edit: I think the iodine has just reacted with the NaOH to form sodium iodate and sodium iodide, the later being easy to separate.
3 I2 + 6 NaOH → NaIO3 + 5 NaI + 3 H2O
I doubt NaIO is that easy to prepare but i cannot find any data on it either way. i will boil down this yellow solution tomorrow and once ive bought more acetone i will attempt to recover as much iodide salts as i can.

[Edited on 20-7-2017 by Assured Fish]
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[*] posted on 19-7-2017 at 22:26


Chemplayer used ammonium peroxydisulfate. The K salt works as well.very fast Available as pool non chlorine oxidisers and they work without acid.
https://m.youtube.com/watch?v=lrwa5x5bTCE
Hate to see others have a lot of trouble then I use this method with peroxymonopersulfate no mixing acid and oxidizer just add and wait works on chlorine bromine and iodine from their respective salts there is chlorine

I once used the same method that you used so I know how frustrating it can be sometimes.

Just add pool chlorine non shock also known as oxone you will be able to retrieve all of the iodine some chlorine will be produced only after all of the iodine is out of solution because of the chloride ion present so you can still salvage the iodine crystals


Also if the was no iodate formed
0.24 g/100 mL (20 °C) for calcium iodate
Side note and off topic
Ca2(IO3)2CrO4 interesting double salt


[Edited on 20-7-2017 by symboom]




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[*] posted on 19-7-2017 at 22:32


@symboom I thought i had watched all of chemplayers vids, guess i missed one, thank you i will definitely opt for a persulfate oxidizer in the future. :D
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[*] posted on 20-7-2017 at 00:54


Oh okay, it was just my guess, because hot, acidic chlorate solutions can oxidize iodine to iodate. And don't be too upset about the loss, I wasted A LOT of iodine in several reactions. One I wanted to recycle like 150g of iodate and periodate waste by adding sulfite to it, but I used a shitty glass container and it broke from the heat of the reaction and I just lost everything. But shit happens, at least you have learned something from this ^-^
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[*] posted on 20-7-2017 at 23:15


ICl in aqueous solution? how is that even possible? one thing is for pretty much sure, that it is not going too be stable, if it is possible whatsoever. I've added (neat) ICl to H2O, a few drops to test its speed of hydrolysis and it was more or less straight away. I really cannot for the life of me envisage it staying around in aq. solution. it decomposes faster than integrity in the presence of politicians.
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[*] posted on 21-7-2017 at 19:56



Quote:

ICl in aqueous solution? how is that even possible? one thing is for pretty much sure, that it is not going too be stable, if it is possible whatsoever. I've added (neat) ICl to H2O, a few drops to test its speed of hydrolysis and it was more or less straight away. I really cannot for the life of me envisage it staying around in aq. solution. it decomposes faster than integrity in the presence of politicians.

You are quite correct, it is not stable, it slowly decomposed producing a orange plume and upon agitation it would decompose quite rapidly.
From what ive read of the other thread which i will reference again, 2 possible answers were given by byko3y and S.C. Wack.
S.C. Wack's post:
Quote:

Brauer, preparation of KICl4: Concentrated KI solution is acidified with hydrochloric acid and chlorine is introduced. The weight increase should be controlled so as to avoid an excess of chlorine. The yield is 70%.

byko3y's post's:
Quote:

Okay, I think now I know what actually happenned. https://www.youtube.com/watch?v=ZFP7fYGrFRI - here's your procedure, http://www.sciencemadness.org/talk/viewthread.php?tid=24862 - and here's a precise description of the result you had and a possible solution to it.

Quote:

I told you! It looks like you've ignored my link, while it was the correct answer - you've got iodine monochloride via I2 + Cl2 -> ICl. Iodine monochloride is somewhat soluble in water, while better also soluble in ethyl acetate. In water iodine mnochloride disproportianates ICl + H2O -> HCl + HOI. That's why bicarbonate is needed - to neutralize HCl (bicarbonate does not form hypochlorite-hypoiodite, only carbonate does). However, I'm not sure whether this reaction is correct or if there's some tricks in measuring reagent amounts..


Truth be told i have no idea who is correct, there are a few threads that talk about this i would advise reading them.
If you want to investigate further here are all the threads I know about, all i know is how it behaved upon baseification and agitation, I managed to form the solution in the first place by using waaaay too much concentrated H2O2 however using dilute H2O2 did not yield the same results.

Please do NOT take my word for it, If you want to know more you are gonna have to do some digging yourself.

http://www.sciencemadness.org/talk/viewthread.php?tid=65425
http://www.sciencemadness.org/talk/viewthread.php?tid=24862
http://www.sciencemadness.org/talk/viewthread.php?tid=5119
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[*] posted on 28-7-2017 at 00:07


Interesting discussion. I have a quart or so of iodine-containing waste, both organic and aqueous. I really have no idea what's in there anymore, other than iodine. I figured the best thing to do would be to add a bunch of NaOH (which I already did), then assume that virtually all of the iodine would be in either the iodide or iodate form. Evaporate all the solvents, (maybe burn some of the organics to see if any purple vapors form) then toss whatever won't dissolve in water. Maybe use ascorbic acid to reduce iodate to iodide (seems like it should work, but never actually tried it), then recrystallize from acetone?



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