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Author: Subject: Formation of polyhalide salt KICl4
woelen
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[*] posted on 12-1-2006 at 14:09
Formation of polyhalide salt KICl4


I did a simple experiment, by adding KIO3 to concentrated HCl and I got very interesting results. After some studies I found that the polyhalide salt KICl4 is formed:

http://woelen.homescience.net/science/chem/exps/KIO3+HCl/index.html

Such polyhalide salts are nice curiosities with really interesting properties. Unfortunately they are not stable. The yellow solid has a strong smell of chlorine, even after one day of leaving it exposed to air. But nevertheless, I thought this is sufficient interesting to post here, the more because I did not yet find a thread about polyhalide salts.

Edit: Made links working again.

[Edited on 26-2-16 by woelen]




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The_Davster
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[*] posted on 12-1-2006 at 16:16


Very Cool

I may well try this later this evening if I have a chance. So how long do you think this salt can store if well dried? Have you attempted to keep any around to see?

This compound brings back the bad memory of on a midterm exam having to draw the lewis structure for the ICl4- ion, and when couting electrons I wrote on my paper "7x5=28" That cost me 20% on the exam(lost marks for carrying through the mistake)....I still got A+ in the course, leading me to believe some of the stuff JohnWW was saying a while back about grade inflation...:P I should bring some of this to show up that proff;)

EDIT: Tried it, very nice:)

[Edited on 13-1-2006 by rogue chemist]




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BromicAcid
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[*] posted on 12-1-2006 at 19:25


Vaguely reminds me of cesium triiodide being stable. The salt of the I<sub>3</sub><sup>-</sup> anion and not with cesium in the +3 state of course. Very interesting, I've never even heard of it....



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[*] posted on 13-1-2006 at 04:41


I've made tetramethylammonium and tetraethylammonium polyiodide salts in the past (during my undergraduate years), specifically the tri-, penta-, and heptaiodides. :) It was messy but fun. The crystals had a metallic look to them. Too bad I forgot to take pictures. :(

Bromic, IIRC, cesium and tetraalkylammonium salts of polyiodides are stable only because they're big enough to form a proper lattice with their counterions.

sparky (~_~)




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[*] posted on 13-1-2006 at 23:12


The crystals shown on your website reminds me incredibly of the crystalline PbI2. The yellow liquid looks also like dissolved PbI2 in hot water. The sharp needle-like crystals is also seen in PbI2 crystals. Perhaps there is similar chemistry here?
Btw, your experiment is a very good curiosity.




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woelen
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[*] posted on 15-1-2006 at 11:53


I also isolated the compound now. It indeed is not very stable. It is not hygroscopic, it forms a nice dry crystalline solid, but this solid emits ICl3 and the crystals become powdered. On longer standing in air, only KCl remains. See the web page I posted before. I have extended this webpage with a lot of text and pictures.

http://woelen.homescience.net/science/chem/exps/KIO3+HCl/index.html

I'm quite sure that isolation of this salt at higher purity should not be difficult at all, if you use a small-volume desiccator with CaCl2 in it, for removing the water. Just air-drying does not work, because then all ICl3 which is given off by the crystals, disappears in the air and only KCl is left.

Edit: Made links working again.

[Edited on 26-2-16 by woelen]




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[*] posted on 3-5-2006 at 10:32
(n-Bu)4N ICl4


Very stabile salt in air and water(insoluble) .Its easy to obtain via precipitation from water solutions of KICl4 with (n-Bu)4NCl and next crystalization from HOT! CH2Cl2 ... that is very good source of ICl4- anions ... ;)




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[*] posted on 3-5-2006 at 12:16


Cool!

For a practicum I've had to make I<sub>2</sub>Cl<sub>6</sub> (or rather (ICl<sub>3</sub>;)<sub>2</sub>;) by letting I<sub>2</sub>, HCl and KClO<sub>3</sub> react. The polyhalide compound that was formed looked a lot like the picture acetate posted.

Nice structure too :) with the Chlorine atoms donating an electron pair to the iodine (each ICl<sub>3</sub> molecule has one Cl donating an e-pair and the I atoms accepting the e-pairs.)



[Edited on Wed/May/2006 by Nerro]
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woelen
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[*] posted on 5-1-2007 at 13:11


I recently made two new polyhalide salts, using CsBr as a starting point. These were really simple to make.

http://woelen.homescience.net/science/chem/compounds/cesium_...
http://woelen.homescience.net/science/chem/compounds/cesium_...

Cesium ion is fantastic in making polyhalide salts. I never imagined it could be so easy. I also noticed that it even can stabilize ions, which are not stable at all in water. As soon as these compounds are dissolved in water, the free bromine resp. bromine chloride are released immediately.

These polyhalide salts make up great oxidizers. The tribromide is a very compact and convenient way of having bromine available. It is sad that Cs is such an expensive metal.

If you have a cesium salt available for experimenting, then it definitely is worth the effort to try to make some polyhalide salt with it.

If you do the same experiments with KBr, then it simply does not work. It does not react with bromine, and if chlorine is passed over it, then KCl is formed and BrCl is released.

If you make such polyhalide salts, be careful not to get them on your skin, they are quite nasty. I noticed that with the CsBr3. It is almost as nasty as pure bromine. As soon as it becomes wet (e.g. by sweat at the skin), it releases its bromine.

Edit: Made links working again.

[Edited on 26-2-16 by woelen]




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[*] posted on 5-1-2007 at 13:55


Why does Cs+ have this ability to stabilize the polyhalides? Maybe because it is "softer" than the other alkili ions?



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[*] posted on 5-1-2007 at 14:08


This reminds me about the accident that happened to me a few years ago. I was trying to make polyhalides. That time I hadn't yet separated pure KClO3 from matches, so I tossed pure match heads to strong HCl-solution. Nothing interesting happened so I tried heating the mixture. I heard a very lound bang. For my luck the test tube was intact but the stuff in it had been blown away. I am not sure where it went and after a quick search I didn't find it.

How very, very stupid I were :P

[Edited on 6-1-2007 by kaviaari]




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[*] posted on 5-1-2007 at 16:20


Quote:
Originally posted by guy
Why does Cs+ have this ability to stabilize the polyhalides? Maybe because it is "softer" than the other alkili ions?

I think it has something to do with the size of the Cs-ion. The bigger, the more stable.




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[*] posted on 6-1-2007 at 04:33


Yes, it is size which matters. Cs(+) ion in fact is a very hard ion, it is most ionic of all ions.

Cs(+) is large, and polyhalide ions also are large, this makes the crystal lattice 'fit' better. Polyhalides of large quaternary ammonium ions (see post of acetate) can also be prepared easily, but unfortunately I don't have such compounds, so I stick to cesium.


Quote:

How very, very stupid I were

Well, I also did something stupid with the CsBr3 I made. I mixed some of it with red P, attempting to make a really cool flame color when the mix is ignited (Br2 + P gives a very violent reaction). But when I mixed the chems, the mix started fuming within seconds :o . My heart skipped a few beats when I saw that, and I immediately dumped it in water, otherwise it probably would have gone KABOOM right in my face. Fortunately nothing bad happened, but mixing polyhalides with red P apparently is not the wisest thing to do! I did not expect the mix to be stable, but that it would be that unstable I did not expect.

[Edited on 6-1-07 by woelen]




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[*] posted on 6-1-2007 at 14:06


Quote:
Originally posted by woelen
Yes, it is size which matters. Cs(+) ion in fact is a very hard ion, it is most ionic of all ions.

Don't you mean soft? Large and polarziable is a characteristic of a soft ion. The soft ions can stablize soft anions like a polyhalide ion.




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