ave369 - 24-11-2016 at 11:22
I've just found that ICl is available OTC as a solution in concentrated hydrochloric acid in my country. It is used as an industrial-strength
veterinary disinfectant. It's quite cheap, it costs an equivalent of $10 for a three liter jerrycan.
What interesting projects can one do with this reagent?
j_sum1 - 24-11-2016 at 11:29
If nothing else you could get a decent quantity of iodine at a relatively cheap price.
Cryolite. - 24-11-2016 at 11:31
Iodine monochloride is a very useful electrophillic iodination. It functions as a source of I+, and can directly substitute an iodine on aromatic
rings, while elemental iodine cannot.
woelen - 24-11-2016 at 11:35
Sure that it is ICl and not ICl3? As far as I know, ICl cannot coexist with water. It forms iodine, iodic acid and hydrochloric acid when dilute and
it forms iodine, tetrachloric iodic(III) acid and hydrochloric acid at very low pH (e.g. in conc. HCl).
The compound ICl3 indeed can be dissolved in conc. HCl and does not hydrolyse. It forms the deep yellow HICl4 (tetrachloro iodic(III) acid). On
dilution, this hydrolyses to iodine, iodic acid (HIO3) and hydrochloric acid. When a potassium salt is added to such a solution, you get crystals of
potassium tetrachloroiodate(III), KICl4.
mayko - 24-11-2016 at 13:57
I was kicking around the idea of using ICl as a source of I+ in an electrophilic aromatic substitution. I got as far as some test tube experiments but
never was able to confirm that a reaction was taking place, and didn't have a good method for separating products. You might get farther than I did:
https://www.sciencemadness.org/whisper/viewthread.php?tid=57...
Recently I ran across some interesting chemistry involving sulfur-nitrogen polymers with various degrees of electrical conductivity. These were
prepared by reacting tetrasulfur tetranitride with halogens or interhalogen compounds. This is probably at the edge of what's doable in a home lab,
though.
AKHTAR, M., CHIANG, C. K., HEEGER, A. J., MILLIKEN, J., & MAcDIARMID, A. G. (1977). Synthesis of Metallic Polythiazyl Halides from Tetrasulfur
Tetranitride. Inorganic Chemistry, 17(6), 1539–1542.
http://pubs.acs.org/doi/abs/10.1021/ic50184a030
Attachment: Synthesis of metallic polythiazyl halides from tetrasulfur tetranitride.pdf (609kB)
This file has been downloaded 375 times
ave369 - 25-11-2016 at 01:32
They say "Iod odnokhloristy" (which means Iodine monochloride) on the jerrycans. ICl3 would be "Iod trehkhloristy".
[Edited on 25-11-2016 by ave369]
woelen - 25-11-2016 at 02:17
What color does the liquid have?
Next weekend I can do a little test. I'll make some ICl and then add 36% HCl to this to see what happens.
ave369 - 26-11-2016 at 08:31
It's orange. The color of orange juice.
woelen - 26-11-2016 at 11:45
I did an interesting observation. I made a solution of ICl3 (by adding solid I2O5 to conc. HCl) and when bubbling of Cl2 stopped and all of the I2O5
was dissolved (which takes quite some time), then I carefully added sodium sulfite to partially reduce the ICl3. This results in formation of chloride
and I2, which in turn reacts to ICl with excess ICl3 and indeed, this works very fine and I got an orange/yellow solution, quite different from the
bright yellow solution of ICl3.
So, ave369, I learned something new. ICl indeed can exist in conc. HCl. Most likely it forms the complex ion ICl2(-). Free ICl is dark brown, but it
becomes much lighter when additional chlorine atoms or ions are attached to it. Free ICl3 is yellow when very finely divided, but orange when in a
lumped state.
Experiments you can do with this:
- bubble a lot of Cl2 through the liquid, until it is bright yellow and then add a concentrated solution of KCl. This will make beautiful crystals of
KICl4.
- You could try adding a concentrated solution of KCl without adding Cl2 first. You may get crystals of KICl2, but I do not know for sure about that.
If you get such crystals I wonder what color they will have. KICl4 is bright yellow, bright as K2CrO4 but somewhat more golden yellow than lemon
yellow. Maybe KICl2 is orange (if you can isolate that).
- Using a carefully controlled amount of Na2SO3 or NaHSO3 you can make solid iodine, which separates from the liquid. ICl is a strong oxidizer, which
oxidizes sulfite/sulphur dioxide and itself it is converted to I2 and Cl(-) anions. Do not use too much sulfite, because that will oxidize the iodine
to iodide and then the iodide dissolves excess iodine to the deep brown I3(-) ion.
Here follows a picture of free ICl, both the liquid and the vapor, which resemble bromine, but are more brown and not as reddish as bromine:
Very interesting that you can buy such reagents OTC!
ave369 - 28-11-2016 at 06:33
The last idea seems the best for me. Elemental iodine is always in short supply in my lab, and making it from medical iodine tincture isn't very
cost-effective. Can I use sodium thiosulphate as reducer? I have plenty of this but little of the sulfites.
[Edited on 28-11-2016 by ave369]
woelen - 28-11-2016 at 12:02
No, thiosulfate decomposes at the very low pH and gives very finely divided solid sulphur particles which will contaminate your iodine and will be
ridiculously difficult to separate from the iodine.
Sulfites are cheap, I expect you can buy them at brewery shops for just a few dollars per kilo. Most common are sodium metabisulfite, Na2S2O5 or
potassium metabisulfite, K2S2O5, which on acidification give SO2:
K2S2O5 + 2H(+) --> H2O + 2 SO2 + 2K(+)
ave369 - 28-11-2016 at 12:27
What about iron (II) sulfate? It's common in gardening shops where I live. For brewery shops, I am not even sure they exist in my country in any form
other than Internet markets.
Cryolite. - 28-11-2016 at 12:49
Can you get sulfur powder? If so, try burning it and bubbling the produced SO2 through a sodium hydroxide solution. This should produce sodium
sulfite, which will suffice as well.
ave369 - 2-12-2016 at 13:31
Yes, I've got loads and loads of sulfur. But constructing the device that traps all SO2 fumes from burning sulfur is kinda beyond me. It needs some
form of positive pressure air pump, and I don't have that.
TheMrbunGee - 2-12-2016 at 14:00
I see that it is only 3% solution, so you have ~90g of ICl for 10 $. Not super cheap, but good for some experiments, I guess! 
CuReUS - 3-12-2016 at 05:16
why not try using sodium dithionite as the reducing agent ? or H2O2 in alkaline medium ?
[Edited on 4-12-2016 by CuReUS]
AJKOER - 3-12-2016 at 08:05
I would add to your ICl/HCl (residing in a wide mouth open vessel), assuming the presence of the ICl2(-) complex (for example, but not necessarily, in
my opinion for subsequent reactions) some N2O dissolved in water. Treat with strong solar light, which would create short lived but powerful hydroxyl
radicals, .OH, that could further react as follows (note: speculation as the direct action of hydroxyl radical on ICl2(-) not documented):
.OH + ICl2(-) =?= OH- + .ICl2
Or:
.OH + ICl2(-) =?= OH- + .I + Cl2
Or:
.OH + ICl2(-) =?= OH- + .Cl + ICl
.OH + Cl- = OH- + .Cl (see https://www.jstor.org/stable/3571859?seq=1#page_scan_tab_con... )
.Cl + .Cl = Cl2
HCl + OH- = H2O + Cl-
..........
In other words, the addition of the hydroxyl radical should generate many reaction chains, with likely a color change and possible liberation of free
iodine and chlorine
[Edited on 4-12-2016 by AJKOER]