Monoamine
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Question: Storing chlorine in solution
I'm wondering if there is a way to make a stable solution of chlorine.
For instance would chlorine dissolve in other chlorinated solvents like DCM or or chloroform, etc..
the reason I'm interested in this is because for Friedel-Crafts mono-chlorinations of aromatic rings.
If you had a, say, 1N solution of Cl2 in DCM (or maybe in Nitromethane since it is slightly polar?, then it would be much easier to
calculate stoichiometry and avoid the problem of over-chlorinating your substrate. If you just bubble in chlorine, then it's very difficult to gauge
if you used enough or if you overshot it.
On the other hand, since adding Cl to the ring deactivates it somewhat, might doing the chlorination at a very low temperature reduce the risk of
double chlorination. (By the way, the compound I have in mind is benzyl alcohol, since it is cheap and somewhat sterically hindered so that one would
expect substitution at the para position mostly I think?
Maybe the answer is using a large excess of benzyl alcohol? Benzyl alcohol is a liquid at room temperature, while 4-chlorobenzylalcohol 68-71 C
(according to Sigma Aldrich). So it may be possible to "freeze out the" 4-chlorobenzylalcohol?
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clearly_not_atara
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Won't chlorine oxidize benzyl alcohol?
Bromine and iodine can be stabilized as Br3(-1) and I3(-1), but the analogous trichloride anion is never stable. The anion BrCl2(-1) is easily
generated by dissolving bromates in HCl, but this is a Br+ equivalent, not a Cl+ equivalent. There is a stable dichlorine complex of
tetramethylammonium hexachlorostannate, (NMe4)4*(SnCl6)2*Cl2, but I doubt this is a practical choice for your situation.
https://pubs.acs.org/doi/abs/10.1021/acs.inorgchem.1c00436
However, hypochlorite in acetic acid can ring-chlorinate activated arenes:
https://www.sciencedirect.com/science/article/abs/pii/S00404...
So I guess that's kind of close?
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j_sum1
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It can be stored in carbon tetrachloride if that helps.
Not 100% sure about chloroform, but it is worth a try.
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woelen
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I made chlorine element samples with saturated solutions of Cl2 in CCl4 and these are stable. I already have these samples for a few years and the
chlorine remains good, even in sunlight.
With CHCl3 I expect more problems. The last hydrogen atom almost certainly will lead to formation of HCl. For immediate use or short term storage,
this may be suitable, but storing Cl2, dissolved in CHCl3 most likely will not work.
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Keras
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I suppose you can keep chlorine in nitrobenzene, too. The nitro- group is sufficiently deactivating to keep the chlorine away from the benzene ring,
at least at r.t.
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Fery
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here solubilities of Cl2 in various solvents:
http://chemister.ru/Database/properties-en.php?dbid=1&id...
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AJKOER
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Interestingly, perhaps one may even consider aqueous based path to "storing" chlorine in high purity water (for example, no Fe or Cu ions) in dark
conditions at a low pH with added chloride.
Logic follows from maintaining a stable HOCl (that is, limit its decomposition pathways) along with moving the below the equilibrium reaction to the
left:
Cl2 + H2O <--> H+ + Cl- + HOCl
As an example, to quote from a 2007 article: "Hypochlorous acid as a potential wound care agent: part I. Stabilized hypochlorous acid: a component of
the inorganic armamentarium of innate immunity", link: https://pubmed.ncbi.nlm.nih.gov/17492050/ :
"Stabilized HOCl is in the form of a physiologically balanced solution in 0.9% saline at a pH range of 3.5 to 4.0. Chlorine species distribution in
solution is a function of pH. In aqueous solution, HOCl is the predominant species at the pH range of 3 to 6. At pH values less than 3.5, the solution
exists as a mixture of chlorine in aqueous phase, chlorine gas, trichloride (Cl(3) (-)), and HOCl. At pH greater than 5.5, sodium hypochlorite (NaOCl)
starts to form and becomes the predominant species in the alkaline pH. To maintain HOCl solution in a stable form, maximize its antimicrobial
activities, and minimize undesirable side products, the pH must be maintained at 3.5 to 5."
by apparently lowering the pH below 3.5 (like from adding concentrated HCl) and increasing Cl- ion presence (with say CaCl2).
[Edited on 21-2-2022 by AJKOER]
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Fery
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In water at low temp Cl2 even forms solid hydrated form Cl2 . 6 H2O
http://www.researchtrends.net/tia/article_pdf.asp?in=0&v...
Attachment: article_pdf.pdf (598kB)
This file has been downloaded 298 times
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teodor
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Quote: Originally posted by woelen  | I made chlorine element samples with saturated solutions of Cl2 in CCl4 and these are stable. I already have these samples for a few years and the
chlorine remains good, even in sunlight.
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Those Cl2 in CCl4 samples are very interesting indeed. Did you make them by bubbling Cl2 into CCl4?
I am wondering whether it is possible to saturate CCl4 by slowly producing chlorine at the same flask which contains CCl4 and then separate the
layers? In this way, one can eliminate the need to set up/clean the gas generator and all that stuff.
[Edited on 21-2-2022 by teodor]
[Edited on 21-2-2022 by teodor]
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woelen
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I made these indeed by bubbling Cl2 in CCl4 (very slowly). I dried the Cl2 with a mix of CaCl2 and P4O10, before dissolving it in the CCl4. For this
reason it is not advisable to use a one-pot reaction with multiple layers. The Cl2 must be dry, otherwise you'll get some water in the mix, which will
react with Cl2 in the long run, producing O2 and HCl and leading to pressure buildup.
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Monoamine
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Thank you for all the advice! I also read that sulfuryl chloride can be used as a source of Cl-, as well as a lot of other cool stuff like radical
chlorination of alkanes, etc...
Made one attempt at making it but the yield was pretty poor. Might try it again though since it seems like a very useful reagent.
Something I haven't tried but learned about recently is something called chloro phthalimide, although the reagents involved to make it seem a bit
precious...
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Texium
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Thread Moved
14-3-2022 at 11:10 |
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