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kmno4
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[*] posted on 22-5-2008 at 03:51


Quote:
Originally posted by kmno4
Using NaBrO3 in oxidation of ethers is interesting in itself.
I am wondering about much cheaper K/NaClO3 as replacement for bromate.


At last I have found some time to check similar rection (THF is the next in queue) with "oxidative dimerisation" of alcohols into "dimeric" esters (as described in given paper). It turns out there is no need to use any compounds containing bromine.
I used KClO3, NaCl in large excess, "catalytic" amount of H2SO4 and n-BuOH. Reaction is strongly temperature dependent, below ~35°C it is very slow and in range 40-45°C it requires about 24 hours to complete (yield >70% of n-butyl butanoate). There are many parameters in this reaction which can be changing and I do not know what set would be the best (concentrations, proportions, temperature, time). There is no generation of gaseous Cl2 or O2 (but organic layer is yellow-green), so I carried out my experiments in closed flask.
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[*] posted on 3-6-2008 at 04:40


Here I am again....
Recently I have found very interesting paper:
The Oxidation of Alcohols by Bromine in the Presence of Bromate
(ja01176a068 , ACS)
Procedure is very nice: single hours needed, very small total volume of reagents, easy workup.
I just replaced KBrO3 with KClO3 (proportionally) and HBr with KBr/H2SO4 (50%).
For oxidation 0,2 mole of alcohol (into ester) 1g KBr is enough.
Reaction temperature should be higher (at the begining), than for KBrO3. Paper says about 50 °C but at this temperature reaction is very slow. I heated mixture at 60 °C (water bath) and after hours it should be slowly increased up to 70 °C. Above 80 °C quick decomposition of HClO3 take place* (ClO2 smell). Reaction is strongy exothemic: temperature must not be increase 60-65°C at the begining because temp. inside flask can be much higher (and dececompose HClO3). Heating at ~70 -75 °C should be continued to a stage when mixture becomes coroless or almost as it. Progress of reaction can be also "monitored" in a simple way: when cooled to ~30 °C, some amounts of salts precipitate but as reaction is more completed, these amouts are become smaller and smaller (c.a ~0,1 g on the bottom of flask in the end). Refluxing is not required (at least at 0,2 mole of alcohol): when solution is colorless, there is no smell of Br2 or Cl2. During experiment flask was stopped and pressure released 2 - 3 times (especially when heated from room temp. → 60 °C ). Upper layer is separated, equal amout of water added and solid NaHCO3 (till no reaction)+ small amout of Na2SO3 (to remove traces of Br/Cl). Organic, upper layer is separated and washed several times with water (water from washings has strongly smell of alcohol).
I used 0,2 mole of n-BuOH as substrate and remaining alcohol is easy to separate from ester in this way. Yield of dried ester (over MgSO4) is 11g. It is about 80% (I did it two times, amounts very similar)) . Please read above-cited paper for more info.

I also tried to oxidate THF in this manner.
Conditions the same, but temperature must be carefully controled, because of boiling point of THF. Using some refluxing aparaturus seems to be safer way in this case (but I also used stopped flask :) ). After reaction is almost [this time I had no time for waiting] complete ( homogenous, light orange color, small amout of salt after cooling), to mixture was added NaHCO3 and Na2SO3. Precipitated salts were filtered and clear, colorless solution obtained ( ~25 cm3). This was extracted 3x3 ml of ethyl acetate. Amounts of product after evaporation, accordingly: 1,5g; 0,7g; 0,3g. Total amount of lactone: 2,5 g from 5,0 g of THF.
Amount of KClO3 was the same as in case n-BuOH, but stoichiometric amount of THF should have been 6,6g. I do not know why, but I used 5g :mad:.
In both cases, stoichiometric amouts of oxidiser are propably a little too small, because of partial oxidation of substrates to corresponding acids.
There exist paper ( CATALYTIC OXIDATIONS IN AQUEOUS SOLUTIONS. 11. THE
OXIDATION OF PRIMARY ALCOHOLS, from ACS) about use of NaClO3 (+H2SO4+ cat. V2O5) for oxidation of alcohols but yields are ~50%
*in this moment orange-brown color of Br2 disappears :o
Propably it is quickly converted by ClO2 to bromate. In this case additional portion of KBr is needed.
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stateofhack
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[*] posted on 18-6-2008 at 04:26


This was not done by me, but from a close relative and i thought that it would be interesting to share and it is to my knowledge the first pictorial of this reaction:

edit: pictures removed because of too much fuss.

[Edited on 19-6-2008 by stateofhack]
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Nicodem
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[*] posted on 18-6-2008 at 08:37


Assuming "sodium GHB" stands for sodium 4-hydroxybutyrate, what makes you believe this is what you obtained, and more particularly what makes you believe what you got is only that. Since some people tend to ingest such a thing (though I never understood why they bother giving that this results in an inebriation quite similar to ethanol), you should understand that such claims could mislead people to poison themselves believing that such a crude product is actually pure sodium 4-hydroxybutyrate while it can contain a number of potentially toxic compounds. Some analytic data would be in order for such claims.



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[*] posted on 18-6-2008 at 09:29
Claret


Having done a lot of organic synthesis and seen the kind of crud that can be found in raw reaction mixtures I doubt that the stuff is anywhere near pure.
At the minimum I would try and get a spot on melting point and TLC and in a real lab, NMR and elemental analysis.

Pass the claret, or a nice ale will do at a pinch...
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[*] posted on 18-6-2008 at 09:30


Quote:
Originally posted by Nicodem
Assuming "sodium GHB" stands for sodium 4-hydroxybutyrate, what makes you believe this is what you obtained, and more particularly what makes you believe what you got is only that. Since some people tend to ingest such a thing (though I never understood why they bother giving that this results in an inebriation quite similar to ethanol), you should understand that such claims could mislead people to poison themselves believing that such a crude product is actually pure sodium 4-hydroxybutyrate while it can contain a number of potentially toxic compounds. Some analytic data would be in order for such claims.


I couldn't agree more, but i am not really hiding anything, i myself really did not do this! :(
My THF has better uses then for such a substance :P
and yes the quality of 4-hydroxybutyrate is dubious indeed, there are without a shadow of a doubt a certain numbers of impurities present in the final product.

I am not encouraging anyone to do something like this! 4-hydroxybutyrate and γ-butyrolactone are illegal in most countries. I just thought it would be nice to have a visual idea of what is going on, picture is worth 1000 words!

I will try to ask the author to get a melting point test :) Does anyone know any other test which could prove that this is indeed 4-hydroxybutyrate?

Althought i suspect the main impurities to be: MeOH, CaCl, NaCl and possible CaOCl.
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[*] posted on 18-6-2008 at 09:47


It is important that you are precise, the impurities probably are MeOH, CaCl2, NaCl and Ca(OCl)2.



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stateofhack
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[*] posted on 18-6-2008 at 10:35


Quote:
Originally posted by woelen
It is important that you are precise, the impurities probably are MeOH, CaCl2, NaCl and Ca(OCl)2.


That is what i said in my post above :P

I do not know how to make this more clear: I did not perform this reaction and so none of these pictures are mine! I do not wish to "add" extra text to the work of someone. If everyone feels that these should not be posted, tell me and i will take them down (or a moderator/admin can do it).

Sorry for all the fuss and i hope that it is now clear :D
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[*] posted on 18-6-2008 at 14:38


To be honest, I'd be more worried about other oxidation products, or even chlorinated organics...
How about reconverting it back to the lactone, and careful distillation?




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[*] posted on 18-6-2008 at 15:08


Yes, wouldn't HCl or H2SO4 work for that? Then extract with nonpolar solvent (not too nonpolar... Like DCM or EtOAc?). And THEN distill off the lactone. I'm pretty sure that should give a decently pure product. I could have sworn that I have read about this exact procedure somewhere before.



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[*] posted on 18-6-2008 at 15:21


OK so this is the stuff that has been getting media in australia (and elsewhere i assume) is it? It has a very narrow therapeutic range (if you would call it that!!) and hence very easily leads to overdoes with symptoms ranging from vomiting nausea, blackout, extreme promiscuity etc. Kind of like having 3 bottles of wine.

I always thought years ago 'wow i wonder what the recreation drug market will morph into as time goes on', knowing trends come and go. This is really LAME!! (in my opinion as an old cunt). I kind of feel happier now about the heroine glut oozing out of Afghanistan, (that is derision, i don't truly feel happy about it).

So what do people do scope a spoonful into their gin and tonic? Maybe the dealers ampule it to ensure quality and dose integrity for their customers.

I suspect there is conjecture amoungst the heirachy here at the site regarding what to do about this obvious blatant post regarding making drugs that society has said it wants to control, but why?

This is not the Hive, we are not bees, and someone you know didn't do this or did. That is only a statement not criticism to try to place how others may perhaps view this.

That said of chemical interest is how mildly the ether linkage is cleaved, i'm a little shocked, if there was isolation of the intermediate so one could truely have a better understanding of the efficacy of this part that would be of interest. Does hypochlorite in these conditions cleave other simple ethers, how do the unsymmetrical ethers behave. If there is in fact very little contamination and relatively pure product is obtainable this easily this reaction has many uses in many situations, by generating functionality in situ. In liquid dimethyl ether say at -30C, what would stop you from making methanol.

:):D




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[*] posted on 18-6-2008 at 17:30


Solvent free permanganate oxidation of THF to GBL (Tet letters 42, 2001)

http://www.erowid.org/archive/rhodium/pdf/solventfree.kmno4....

There is no mention of ratios of oxidant etc. in this paper.

Can someone suggest an appropriate amount of oxidant per mole of substrate for this reaction?

I am sure a large excess would be used.
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[*] posted on 19-6-2008 at 03:14


There is enough information for a competent chemist to reproduce the reaction if you can be bothered to read the paper.

If you are not a competent chemist maybe you should ask yourself if you should be trying to turn THF into GBL at home.
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[*] posted on 19-6-2008 at 04:25


Yes.. I have read the paper, thanks for asking ;)

Yes.. the patent contains some useful details on the preparation of the oxidant and the yields of the various oxidations.

Yes.. I am a competent, educated chemistry enthusiast.

No.. The paper does not have an experimental section or does it clearly describe appropriate ratios of oxidant to substrate for the various oxidations.

There is, however, the mention of 2 mmol of substrate to 4 grams of oxidant. This would equate to, by my calculations, approximately 1 mol of THF to 2000 grams of oxidant and would therefore be unfeasible.
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[*] posted on 19-6-2008 at 04:31


"There is, however, the mention of 2 mmol of substrate to 4 grams of oxidant. This would equate to, by my calculations, approximately 1 mol of THF to 2000 grams of oxidant and would therefore be unfeasible."

No.

You would actually need 8Kg (8000g)

2 x 1000 x 4 = 8000
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[*] posted on 19-6-2008 at 08:52


Oops, got a bit over keen there, it is indeed 2kg.

As the method is fairly impractical, I will tell you how I would go about it.
Take 2g of fine Al203 and place it in a mortar, add saturated potassium permanganate solution until a stiffish paste is formed. Then add the copper sulphate powder and grind the mixture until it forms a crumb.
Add to the 25ml flask containing the stirrer bar and seal with a rubber septum. Start stirring and inject 0.144g of THF.
Allow to stir for 6 hrs and then inject 10ml of dichloromethane.
Unstopper the flask and filter at the pump, washing the solids with a further 5ml of dichloromethane. Dry and filter into a weighed flask. Concentrate on the rotavap and weigh again.
That gives you a yield which I think is low in this case as they sucked some of their product into the pump :D
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[*] posted on 19-6-2008 at 15:29


Sweet, maybe after a few days I might have 1 ml.

This would not be enough to clean the ports on my engine.

This reaction would be attempted at no less than 1 mole.

They likely used such a huge excess in paper, as they often do, because they are working on such a small scale and for analytical reasons.

It may well work with 100 grams of oxidant for 1 mole of substrate if stirred properly and allowed to continue for many hours for all we know.

Was hoping to generate some discussion on this point before jumping headlong into it. But I guess this is what experimenting is all about.

Anyone with experience or an opinion, please chime in.
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[*] posted on 19-6-2008 at 15:57


You could do a whole series of experiments, where the 'analytical' step is quantitative distillation... then see at what excess of oxidant the yield starts remaining constant.
Also I guess you could try other strong oxidisers- such as K2Cr2O7 - it is maybe more expensive but maybe a lesser excess is required.

The reaction is interesting generally I think, given the wide variety of compounds that can be oxidised!

The paper Natural points to also mentions microwave, so that's something to consider.

Also, no need for such a tone here (particularly being new to this forum), ScienceSquirrel.




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[*] posted on 20-6-2008 at 04:13


"This would not be enough to clean the ports on my engine."

Assuming that you live in the good old US of A where this stuff is illegal, isn't it a bit hard core to try and make a controlled substance to use as a cleaning agent?
However it does bring alloy wheels up a treat and removes chewing gum with ease but it will take your paint off in a blink of an eye.
It is available in Western Europe by mail order so you could always order a bottle and plead ignorance if the US Customs picked it up...
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[*] posted on 21-6-2008 at 05:56


I'm sure no one is interested in where I'm from or what I do.

There is a follow up article to the one listed above in which they use a different support for the KMnO4 and list the yields for oxidations of various benzylic compounds under microwave conditions etc. The yields are good.

They still use a similar excess of oxidant, but aren't they always doing this in these tetra papers to guarantee max conversion without having to make adjustments?

It will most likely be attempted at a 3:1 ratio, in the microwave on medium for 30 mins as per the more recent paper I have.
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[*] posted on 21-6-2008 at 06:05


ScienceSqirrel, technically it is no more illegal than red phosphorus or iodine. It's just that our legal system doesn't care what is (il)legal anymore. As long as we "catch the badguys".

So, I suppose it is illegal by proxy. Therefore, yes, it does seem a little extreme to use it as a solvent.

Why is it that so much oxidant is required for this? Isn't mixing an oxidizer such as KMnO4 with an ether pretty dangerous?




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[*] posted on 22-6-2008 at 07:07


A lot of oxidant is required as it is only effective on the surface and most of the material is unavailable for reaction.
This type of reaction is fine for making stuff on a sub gram scale in the lab if the reagent is highly efficient or shows some unique specificity.
For making GBL on a preparative lab scale I would use the calcium hypochlorite method
http://www.erowid.org/archive/rhodium/chemistry/ether2ester....
You could probably scale it up a few times with no problems and use just acetic acid / water as the solvent due to the high solubility of THF in aqueous solutions. The 60% acetic available for stop baths etc should be fine.
Your main costs would be a heater stirrer and a good distillation apparatus.
You would want to distill off and reuse your dichloromethane as it would be quite expensive bought on a litre scale.
If you were making 50 - 100ml at a time you would soon have plenty to wash your wheels.
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[*] posted on 13-7-2008 at 18:50


Oxidation systems utilizing H2O2 catalyzed by a group 6 metal ion and a PTC oxidize secondary alcohols to ketones and primary alcohols to carboxylic acids in good yields. There are several papers covering this topic.

IIRC most systems that oxidize primary and secondary alcohols also oxidize ethers to esters and cyclic ethers to lactones effectively.

Would it be a safe assumption that this would also be the case for the H2O2/metal-ion/PTC system?

No examples of this could be found.
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[*] posted on 1-12-2009 at 14:35


Advanced Organic Synthesis METHODS AND TECHNIQUES RICHARD S. MONSON page 12
Oxidation of Ethers to Esters
The oxidation of ethers to esters according to the reaction offers many possibilities
for the modification of functionality in open chain or cyclic systems. An example is the conversion of tetrahydrofurans to y-butyrolactones. Two reagents have been discovered
that allow for this conversion in satisfactory yield: ruthenium tetroxide and
trichloroisocyanuric acid (Chapter 17, Section IV). The use of these reagents is given
below for the conversion of di-n-butyl ether to n-butyl n-butyrate.
n-BuTYL BUTYRATE FROM DI-W-BUTYL ETHER BY TRICHLOROISOCYANURIC ACID (17)
In a 200-ml round-bottom flask equipped with a magnetic stirrer and a thermometer
is placed a mixture of 50 ml of di-n-butyl ether and 25 ml of water. The flask is immersed
in an ice bath and the mixture is cooled to 5°. In one portion is added 23.2 g (0.1 moles)
of trichloroisocyanuric acid (Chapter 17, Section IV), and stirring in the ice bath is
continued for 12 hours. The ice bath is removed and the mixture is stirred at room
temperature for an additional 8 hours. The reaction mixture is then filtered to remove
solids. The water is separated from the organic layer, which is then washed with two
additional portions of water, dried with anhydrous sodium sulfate, filtered, and
fractionated as above.
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[*] posted on 2-12-2009 at 01:59


Quote: Originally posted by The_Natural  
Oxidation systems utilizing H2O2 catalyzed by a group 6 metal ion and a PTC oxidize secondary alcohols to ketones and primary alcohols to carboxylic acids in good yields. There are several papers covering this topic.

IIRC most systems that oxidize primary and secondary alcohols also oxidize ethers to esters and cyclic ethers to lactones effectively.

Would it be a safe assumption that this would also be the case for the H2O2/metal-ion/PTC system?

No examples of this could be found.


I would personally avoid using THF with peroxydes...

I've read a very good report about the oxidation of THF with Ca(OCl)2. The yield is about 60% and the method is totally OTC. I don't want to post it because I hate clandestin drugged chemist... They have no interest about chemistry, they just want their drug.
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