Sciencemadness Discussion Board

Oxone Aromatic Halogenation in Water - Is this write-up a lie?

FireLion3 - 20-7-2014 at 13:38

I know, the title sounds stupid, but I don't really have much of an explanation here. This is my 4th time trying this prep identically to the write up with compounds they used in the write up and have had nothing but failure.

http://www.researchgate.net/publication/237152679_Direct_hal...

First I attempted it with Anisole, following their write up and molar quantities precisely. After several hours of stirring, then letting sit, the organic layer at the bottom was dark red/brown. This was a few weeks ago, and I made a thread curious on whether I had gotten the wrong reagent from my supplier because it did not react as anticipated.

Lacking the equipment to verify if the compound was correct, I purchased some 1,3-dimethoxybenzene from a highly reputable supplier. I repeated the prep with this new reagent on two different occasions, and I continually got the same result. Dark brown/red layer at the bottom no matter how long I stirred.

Now, this prep claims these two compounds, and several electron rich aromatics, should complete reaction in 5 minutes, at room temperature. With each two trials for each compounds, one was run at room temperature, the other at 80 degrees. The electron rich compounds are more prone to electrophilic attack by the bromine electrophile created from Oxone + NaBr. The Oxone and NaBr used were as pure as they get.


These brominated compounds are supposed to be clear. The red color cannot possibly be the target-compound because even with slight stirring it mixes very rapidly into the aqeuous layer, but after stopping stirring it begins to gravitate towards the organic layer - whereas the organic compound itself very strongly stays towards the bottom. So, solubility cannot possible be an issue, since the bromine electrophile appears to favor the organic layer. The only thing that I can possibly think of is my stir rate not being fast enough. My stir plate wont let me go over 700 RPM or it stops stirring, but I can't imagine how this would be a major issue since the bromine favors the organic layer. Especially, not a big enough difference for the authors to get 100% yield in 5-10 minutes at room temperature and have be get no noticeable reaction progress in 2-4 hours at 80 degrees.

Has anyone else attempted this reaction? It is as fairly simple and straightforward as one can get. Why does there appear to be such a huge amount of unreacted bromine? I'm going to add an excess of aromatic in a little while to see if that could possibly be the cause... but I know my digital scales are not *that* inaccurate. There's no reasonable explanation to why there would be a large excess of bromine.

[Edited on 20-7-2014 by FireLion3]

I will attempt to get some pictures up later.

[Edited on 20-7-2014 by FireLion3]

forgottenpassword - 20-7-2014 at 14:30

Your product has dissolved the excess bromine. Wash it with sodium bisulfite solution.

FireLion3 - 20-7-2014 at 14:49

How can there be excess bromine? When I stop stirring, the bromine sinks to the bottom, when I start stirring again, it quickly fills the entire water portion - so the bromine being trapped is out of the question. There is about 1 Liter of water and half a mole of arene (50ml of arene). An hour ago I added an extra molar portion of the arene (so there is one mole now) to react with any excess bromine (there is no way I added 2x excess sodium bromide when starting), and the red color has not reduced at all. Solubility is no explanation since the bromine appears to so easily leave the organic layer.

How did the authors have such great success? I'm debating on switching over to using N-Bromosuccinimide if this can't work.

I have access to an overhead stirrer that will give much better stir rates, but I am currently waiting on a stand for it so until then I cannot use it.

[Edited on 20-7-2014 by FireLion3]

[Edited on 20-7-2014 by FireLion3]

forgottenpassword - 20-7-2014 at 15:11

You've scaled up by 1000 times what was reported in the paper. Don't expect the reaction time to be the same.

FireLion3 - 20-7-2014 at 15:21

While that sort of reasoning is true in many reactions, it should be close to negligible in a reaction such as this where the necessary reactants are all basically in the same phase and readily react. With the strong activating groups on the aromatic ring, there should be practically no inertia for the bromine electrophile to react with the arene. Also consider I am running the reaction 50 degrees higher than required for it to take 5 minutes, it should in theory be running 32x faster. This apparent lack of reactivity given the circumstance cannot so easily be explained away by a mere scale up.

It's still stirring so I am just going to give it time, and if needed, run it 24+ hrs unlike when I was unable to prior.

Tomorrow I should be able to run the reaction again with my overhead stirrer so perhaps that will provide a noticeable enhancement via RPM increase.

forgottenpassword - 20-7-2014 at 15:26

Ok, expect it to take the same amount of time. Why ask questions if you know the answers?

FireLion3 - 20-7-2014 at 16:00

Are you intentionally going out of your way to twist my words and make fallacious statements? Please don't troll me. You posted one line responses without going into any detail or reasoning whatsoever, then you derogatively accuse me of knowing everything because I am questioning your one liners with something called reasoning? If you're unwilling to engage in critical thought or answer questions about your own statements, then why are you on a message board? So please, save your thinly veiled insults and do not patronize me when I ask for clarification. If you don't know just say so.

Your two ideas are reasonable among other circumstances but glancing at this paper would quickly inform that those two factors are unlikely to be the cause here given the specific reaction circumstance - Unless of course you are aware of some unstated mechanism? If so, please elucidate.

I'm not asking a lot by requesting an explanation that relates to this specific reaction that extends beyond a single line.

mnick12 - 20-7-2014 at 16:26

FireLion3, I understand your frustration with what appears to be a failed reaction. That being said I wouldn't take your frustrations on the community you are asking help from.

The thing I noticed in your description is the lack of any sort of work up or attempt at characterizing your reaction product. How can you claim the procedure is failed if you never even worked up the reaction? From the procedure they washed the reaction with Na2S2O4 and extracted with ether. I would suggest you work up you reaction and do some sort of characterization before you draw any conclusions.

As far as brominations go, from my own experience there is always a red color left over. Even if the organics are used in excess.

FireLion3 - 20-7-2014 at 17:38

Thanks for the reply. I wasn't attempting to take my frustration out on the community, just became rather irritated when I was insulted for asking him to clarify his answer.

Here is the thing about this reaction: In the paper they claim quantitative conversions, and they did not use any excess Oxone or NaBr. This would imply that there would be no excess, or very little, bromine remaining. The less than 100% isolated yield can be explained by slight solubility of the organics in water. If I was using something less reactive like benzene I would expect the workup..... but this paper, and several other papers, have claimed 100% conversations with electron rich arenes so ideally the workup could be restricted to a simple phase separation - inevitably some would be lost in the water if not extracted though. With a 100% conversion there's no explanation for the extremely dark red color, unless they of course lied about that figure.

I wasn't planning any workup until the red color was nearly completely gone, perhaps with some slight yellow remaining. Though, after 6+ hours of stirring and heating, the solution is just as red as it was when I added the last bit of the oxone. They used Sodium dithionite to remove residual bromine from the organic, which in most cases there was none if there was a true 100% conversion, and then they extracted the aq. with ether and dried it with Sodium Sulfate.

Question for you:
Did you perform your brominations via this method, and if so on what compounds?

Edit:

One thing I am unsure of is how Bromine even forms in the first place. Bromine is in equilibrium with Hypobromous acid, but it requires Hydrobromic acid to turn into bromine, which is not present at all in the solution. Oxone converts the bromide ion directly into hypobromous acid, which should then turn into water after exchanging its bromine for one of the protons on the aromatic ring.


[Edited on 21-7-2014 by FireLion3]

mnick12 - 20-7-2014 at 19:47

Well conversion is different than yield, as you seem to point out in your post. Again you cannot talk about yield until you do some sort of work up. Perhaps you may not get the 100% yield you hoped for, but you must do something before you make such claims.

I have never done brominations with oxone and bromide, however I have done plenty with bromide and peroxide and even elemental bromine. In both cases the product is usually orange to red and requires at least one re-crystallization before the mp is acceptable. I have used these brominating techniques on vanillin, 1,4-dimethoxybenzene, and ethylvanillin. I would follow the work up described and see what your yield is.

FireLion3 - 20-7-2014 at 20:23

Quote: Originally posted by mnick12  
Well conversion is different than yield, as you seem to point out in your post. Again you cannot talk about yield until you do some sort of work up. Perhaps you may not get the 100% yield you hoped for, but you must do something before you make such claims.

I have never done brominations with oxone and bromide, however I have done plenty with bromide and peroxide and even elemental bromine. In both cases the product is usually orange to red and requires at least one re-crystallization before the mp is acceptable. I have used these brominating techniques on vanillin, 1,4-dimethoxybenzene, and ethylvanillin. I would follow the work up described and see what your yield is.


Does Peroxide + Bromide work? I was under the impression that Hydrobromic acid was required for peroxide.

In my case I do not think recrystalization will be so easy since these are predominantly liquid around room temperature.

Since I don't have anything around to clean this excess bromine out, I went ahead and added 300mL of Toluene. Toluene is also supposed to brominate very rapidly. So, in solution now there is nearly a 1:6 excess of Bromine/Hypobromous to Aromatic. According to my timer it's been roughly 8 hours since I've started the reaction. I will let it go overnight and see what happens in the morning.

If there is no reaction overnight then I will assume that my stirring is the problem. With this much excess arene, all of the bromine should be consumed - eventually - but it really should not take that long seeing as even toluene reacts extremely readily with the bromine electrophile.

forgottenpassword - 21-7-2014 at 04:00

Judging by the fact that you have lots of unreacted bromine floating about it should be obvious that the reaction is not instantaneous. Scaling up a procedure by 1000 times will always take longer than the 1/1000 scale procedure. In the paper they talk of 20/1000 of a mole as being large scale. Using a solvent that dissolves your starting material would make more sense on such a large scale as you are using.

As mnick said, you can say nothing of the yield until you have purified your product. Now that you have contaminated it with bromo-toluenes that will be much more difficult. A simple wash to remove excess bromine, as they employ in the paper, would make much more sense. They had excess bromine, and you have excess bromine.

[Edited on 21-7-2014 by forgottenpassword]

Nicodem - 21-7-2014 at 06:04

Quote: Originally posted by FireLion3  
but this paper, and several other papers, have claimed 100% conversations with electron rich arenes so ideally the workup could be restricted to a simple phase separation - inevitably some would be lost in the water if not extracted though. With a 100% conversion there's no explanation for the extremely dark red color, unless they of course lied about that figure.

What you say is nonsense.
Do you think products just jump out of reaction mixtures? Utmost they can precipitate and can be isolated by filtration, washings and drying. That's far from jumping out of a reaction mixture, but the closest you can get. Obviously, not applicable to your reactions.
Why do you imply the 100% conversion as having anything to do with "extremely dark red color"? You were brominating aryl methyl ethers. What other color did you expect? Besides, the substrate conversion says nearly nothing about yields. You can have a 100% conversion with 5% reaction yield and 1% isolated yield - so what? In fact, the only thing that substrate conversion says in regard to yields is that the yield cannot be higher than the conversion.

And most importantly, how do you dare accusing the authors of telling lies without bothering to either monitor the reaction or isolate the products? You say that based on what rational argument? You do not even bother comparing their result with other reports.

While I would not trust that article, I would at least check the original articles from which they got the idea. I would also rely the general knowledge that many of the used substrates brominate without troubles with elemental bromine (so they will obviously brominate with a bromide and oxidant combination). Even in lies can there be accidental truth.
Quote:
They used Sodium dithionite to remove residual bromine from the organic, which in most cases there was none if there was a true 100% conversion, and then they extracted the aq. with ether and dried it with Sodium Sulfate.

They used the dithionite to quench and discolor like it is common for electrophilic brominations. There can be no bromine where the substrate is anisole or 1,3-dimethoxybenzene as it would react to give the dibrominated products. Try isolating the products and you will probably find out that the color does not distill over (unless on 1,4-dialkoxybenzenes as substrates where volatile colored compounds tend to form - benzoquinones are colored and volatile).
Quote:
One thing I am unsure of is how Bromine even forms in the first place. Bromine is in equilibrium with Hypobromous acid, but it requires Hydrobromic acid to turn into bromine, which is not present at all in the solution. Oxone converts the bromide ion directly into hypobromous acid, which should then turn into water after exchanging its bromine for one of the protons on the aromatic ring.

Please avoid speculating about things you obviously do not understand. This is not the Beginners section, so cite references for claims that do not make sense.

FireLion3 - 22-7-2014 at 09:55

So I made an interesting discovery today that explains a lot! In the process I found a very effective OTC way to produce anhydrous Chlorine and Bromine gas.

forgottenpassword was partially right. Bromine is very soluble in the organic layer (my reagent). But how did Bromine form I wondered? Hypobromous Acid is in equilibrium with Bromine IF Hydrobromic Acid is present, AKA, Hydrogen Bromide. Hypobromous acid is clear so the red color has to be bromine. So... where did the Hydrobromic acid come from? The answer: There was none, but there was something else...

Last night when experimenting with DBDMH. https://en.wikipedia.org/wiki/DBDMH. In aqueous solutions this disassociates to hypobromous acid, but can also act as a source of radical bromine in some cases.
Quote:

https://en.wikipedia.org/wiki/DBDMH
DBDMH is a source of bromine, which is equivalent to hypobromous acid (HOBr).

Br2X + 2 H2O → 2 HOBr + H2X


I had DBDMH stirring with an aromatic substrate containing no methyl side chains, and for 30 minutes there was no observed color. Out of curiosity I decided to add some Sodium Bromide. Immediately the solution turned red. Also, when adding the Sodium Bromide through my funnel, it came into contact with some of the dry DBDMH, and it instantly turned red and some red gas evolved. The more bromide I added to solution, the redder it got. So there is one explanation here. Hypobromous acid is in equilibrium with Sodium Bromide to form Bromine and Sodium Hydroxide.

So instead of:
Quote:

https://en.wikipedia.org/wiki/Hypobromous_acid
When pure bromine is added to water, it forms hypobromous acid and hydrobromic acid (HBr):

Br2(l) + H2O(l) ↔ HOBr(aq) + HBr (aq)

We have:
Quote:

Br2(l) + NaOH(l) ↔ HOBr(aq) + NaBr (aq)

Or alternatively (X = DBDMH/BCDMH):
Br2 + NaX ↔ XBr + NaBr

If we take a look at a similar compound:
https://en.wikipedia.org/wiki/Sodium_dichloroisocyanurate
https://en.wikipedia.org/wiki/Cyanuric_acid

We can see that with Cyanuric acid the Tautomers interconvert. Sodium Hydroxide can be reacted with this compound to form the sodium salt. I am not sure nor do I have any references for it, but I believe the compound can be deprotonated up to 3 times.

So basically with the second equation above (in this quote box), it may be broken down in solution as (An enormous equilibrium process):

1) XBr(aq) + H2O -> HOBr(aq) + X-H
2) HOBr(aq) + NaBr(aq) -> Br2(l) + NaOH(aq)
3) NaOH(aq) + X-H -> NaX(aq) + H2O

Consequently we also have:

4) Br2(l) + H2O(l) ↔ HOBr(aq) + HBr (aq)
5) HBr(aq) + NaX -> X-H + NaBr
6) NaBr + HOBr + NaOH + Br2




To test this theory, I repeated it several more times. Each time, only when I added bromide to DBDMH, dry or an aqueous solution, bromine gas evolved. I even took some Trichlorocyanuric Acid and added Sodium Chloride to it, and Chlorine gas evolved! In a separate case, I added bromide to Trichlorocyanuric acid and Bromine Gas formed! This latter case makes sense seeing as when Hypochlorous acid reacts with Bromide, Hypobromous acid and sodium chloride are formed. The hypobromous acid does not react with the sodium chloride, but will react with bromide anions.

This is interesting because when I was shopping I noted that I passed a container full of Trichlorocyanuric Acid that had a warning on it saying do not mixing with any other products, sodium chloride, hypochlorous acid, sodium bromide, etc. I see why now! Dangerous gas evolves!
Quote:

https://en.wikipedia.org/wiki/BCDMH
The bromide ions are oxidized with the hypochlorous acid that was formed from the initial BCDMH:

Br- + HOCl → HOBr + Cl-

This produces more hypobromous acid. However, the hypochlorous acid itself does act directly as a disinfectant in the process.


I don't know why, but I was assuming Hydrogen Bromide is different from Sodium Bromide, but they both contain the Bromide ion. Before I get on to quoting Nicodem, I will thank Nicodem for a post of his I read of his when he stated that he was able to produce bromine from Hydrogen Peroxide and Sodium Bromide. If you google search brominations with hydrogen peroxide, they almost all exclusively discuss using Hydrogen Peroxide and Hydrobromic acid.

The above that I have discovered is very interesting seeing as I haven't found it in any papers. Whatever the case may be I WAS able to successfully brominate my compounds with just DBDMH in water No bromine gas formed, or any other red color. At the end of the reaction, while I cannot qualitatively determine products, the organic layer weighed much heavier than initially, indicating bromination, and the aqeous layer was also much heavier. DBDMH is not very soluble in water, but it's dehalogenated form that results after it gives up its bromine "5,5-Dimethylhydantoin", IS soluble in water. I found reports of a similar reaction with these compounds being done with NBS that claimed to be finished in 5-10 minutes at RT, but I let my run at 80 degrees for 2 hours.

As a further form of verification, I went on to test the brominated compound and it reacted as expected. I am most thrilled because there were no red impurities that I had to deal with at any point in this process. A few water washes, phase separation, and drying with magnesium sulfate was all I needed.



I will note that I typed this below response to you Nicodem before typing the above, but I put the above first as it directly pertains to the thread, while below I am responding directly to you
Quote: Originally posted by Nicodem  

Why do you imply the 100% conversion as having anything to do with "extremely dark red color"? You were brominating aryl methyl ethers. What other color did you expect?

Many brominated and other ring-halogenated compounds are clear, especially the ones that I was working towards. You don't think I would look up the color of my expected compound before doing a reaction? Also, I would appreciate it if you didn't post without reading, you are making it very apparent that you are doing a lot of that lately. I may be a newbie at chemistry but I am not that big of a fool
Quote:
Besides, the substrate conversion says nearly nothing about yields.

Considering the compounds I was using that the authors also used claimed 100% conversion and 93%+ yields... I would say that the word "yields" is indeed talking about "yields", especially when they display a picture of the monobrominated product right next to the word "yields".
Quote:

There can be no bromine where the substrate is anisole or 1,3-dimethoxybenzene as it would react to give the dibrominated products.

That is correctly, but dibrominated compounds generally only form is stirring is very poor and there is a large excess, OR if there is a lewis acid catalyst to form the more electrophilic bromine intermediate complex. In electron rich compounds, the nonbrominated compounds will obviously brominate much faster than the monobrominated ones, but if there is excess bromine and enough time given then there will be some dibrominated compounds. If a lewis acid catalyst is used, there has been shown to be a much ratio of mono-brominated and di-brominated compounds due to the very high electrophilicity of the lewis acid halogen complex.
Quote:

Please avoid speculating about things you obviously do not understand. This is not the Beginners section, so cite references for claims that do not make sense.

Should I assume that every since piece of common knowledge requires a reference? Well Tally Ho mate since you asked. You claim these things do not make sense but they should be common knowledge for anyone who has ever touched a halogenation reaction, that is assuming you read up on the reactions before you do them? I never touch a reaction without reading as many articles I can find on it, and if you do this you will see little details like this are indeed very basic information. I can quite confidently say there is not a single reaction I have done that I have not read at least 50+ papers on about that specific topic.

Quote:

https://en.wikipedia.org/wiki/Hypobromous_acid
When pure bromine is added to water, it forms hypobromous acid and hydrobromic acid (HBr):

Br2(l) + H2O(l) ↔ HOBr(aq) + HBr (aq)

https://en.wikipedia.org/wiki/Hypochlorous_acid
Addition of chlorine to water gives both hydrochloric acid (HCl) and hypochlorous acid:[8]

Cl2 + H2O is in equilibrium with HClO + HCl



Avia22 - 1-8-2014 at 13:59

FireLion,

I tried to dissolve dbdmh in heated and stirring water, and it did not seem to dissolve. Write ups say it should react and turn into hypobromous acid and DMH. DMH is supposed to be soluble in water, so it seems this reaction isn't occuring.

I added methoxy-benzene to this mix after minimal reaction observed in 1 hour of stirring, and the solution turned very red. I am not sure how this happened. I added an excess of methoxy benzene and let stir for 2 hours at 40 degrees, and there was still lots of red color (total black appearance once let settle). I saved my reaction mix in a plastic bottle, and am waiting on some bisulfite to arrive so I can see how much it takes to remove leftover bromine.

How did you get your reaction to work without generating any red color? I also tried adding the dbdmh directly to the methoxybenzene without water and there was still lots of red color. The MSDS sheets I read said bromo-methoxybenzene is a "pale yellow" color, what I have is very dark red, near black. What am I doing wrong? The dbdmh is obviously reacting but I don't know why the solution is so dark red or why bromine forms. Without water present, shouldn't the dbdmh just undergo substitution with the nucleophile hydrogen on the ring?

FireLion3 - 1-8-2014 at 14:52

Avia I will need to repeat this reaction on a bigger scale. I suspect my not encountering any red color was due to the low concentration I used. I did some reading and found some interesting things.

This:
http://www.arkat-usa.org/get-file/32614/

And attached Below.
Quote:

The N-sulfinyl α-amino-1,3-dithianes 8 were hydrolyzed by treatment with 2 equivalents of DBDMH 4 in 80% acetone at -20 oC. The solution quickly turned red and then faded to yellow-orange after a few minutes. The reaction was quenched after about 10 min by addition of aqueous sodium sulfite to afford the crude N-tosyl aldehydes (S)-9 as colorless oils.
Quote:

Solid DBDMH (0.50-0.52 mole equiv.) was
added in part into the solution of starting
material (Ar-OH) in CHCh (5-7 mIl mmol) at
room temperature. Upon addition of the
DBDMH, the solution became red or deep
brown colored, the next portion of DBDMH was
added after the disappearance of color and so on.


Your description of the solution appearing near-black comes close to the description above it a deep-brown color.

Both documents state that
immediately upon addition of DBDMH the solutions turned red. This leads me to believe that since DBDMH is soluble in organics, when dissolved, it has a natural red color, since the crystals themselves are light-yellow.

In the attached document it seems most of their brominations with electron rich compounds were complete within 2 hours, with some taking 12 hours, yielding 75-98% with some cases having trace amounts of double brominated compounds. These all were run at room temperature so I'm sure the speed could be increased at higher temperatures.

Both articles seem to quench with sulfite's. I think the sulfite might react with the bromine on the DBDMH and neutralize it, forcing it out of the organic layer since DMH is soluble in water.

I don't think the Red Color is bromine itself. It seems to appear very rapidly upon addition of addition of the DBDMH, and there's no plausible reaction for how the bromine could form with only organics in the solution.

I noticed myself the DBDMH doesn't dissolve in water well with heat, but I noticed when it was stirring and heating for 30+ minutes, there was a bubbly/fizzy/sizzly layer building on top of the water that was not there from the start. I believe this was hypobromous acid, since there's not much else it could be. The DBDMH itself never fully dissolved and I believe the reaction that forms hypobromous acid is very slow. There doesn't seem to be much of a point to the water at all since the DBDMH dissolves right into the organic layer.



Attachment: 010_105_109.pdf (196kB)
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EDIT: BACK WITH AN UPDATE:

Avia, I would say redo your experiment expect keep the water out, or make sure the vessel is dry.

I just added Anisole with some (noncrushed) flakes of DBDMH into a spare bottle I had. I shook this and immediately it turned red. The DBDMH did not dissolve immediately since it was not in powder form, but I can tell it is dissolving.

I would shake the bottle, let it sit and over the period of 2 minutes the red color would fade to orange than to yellow. I observed clearly that there was only one single color layer. I shook the bottle again and it instantly turned red, since there were still flakes of DBDMH still dissolving. The reaction definitely seems to be taking place.


[Edited on 1-8-2014 by FireLion3]

FireLion3 - 2-8-2014 at 17:35

[New Post]

This is very strange. The reaction itself seems to take place very rapidly when there is a small amount of brominator added to. The color can be observed going from dark brown to light yellow. However, when a stochiometric amount is mixed, the mixture can be stirred for up to 18+ hours, as I have found out, without any lessening of the dark color. Even heating at 100 degrees for 5 of those hours did not have any effect.

At room temperature in a premixed bottle, sitting still, the reaction takes place fine, but does not seem to progress at all when stirred with stochiometric quantities at RT or with heat. This is with activating groups on the ring.

I have no explanation. I suspect using a strong Lewis Acid like Aluminum Chloride could accelerate this reaction either by complexing with the halogen, or by complexing with the carbonyl group on the ring to make the compound more electrophilic. Does anyone have any suggestions?