Sciencemadness Discussion Board

thf oxidation

runlabrun - 5-1-2005 at 01:04

Just a question about the various oxidation pathways for thf, producing gbl.

Sodium bromate:
Reagent: NaBrO3, NaHSO3
Solvent: H2O
Time: 16.25 hours
Temperature: 20 C
Yield 68%
Ref. 1 6148927; Journal; Sakaguchi, Satoshi; Kikuchi, Daisuke; Ishii, Yasutaka; BCSJA8; Bull.Chem.Soc.Jpn.; EN; 70; 10; 1997; 2561 - 2566;
or -->
Reagent: aqueous sodium bromate
Catalyst: 47% hydrobromic acid
Solvent: CH2Cl2
Time: 5 hours
Temperature: 35 C
Yield 67%
Ref. 1 5934286; Journal; Kajigaeshi, Shoji; Nakagawa, Takashi; Nagasaki, Noritaka; Yamasaki, Hiromochi; Fujisaki, Shizuo; BCSJA8; Bull.Chem.Soc.Jpn.; EN; 59; 3; 1986; 747-750;

But also what about :
Reagent: Mn2O7
Solvent: CCl4, acetone
Temperature: -45 C
Yield 86%
Ref. 1 5689808; Journal; Troemel, Martin; Russ, Manuel; ANCEAD; Angew.Chem.; GE; 99; 10; 1987; 1037-1038;
or -->
Reagent: ZnCr2O7*3H2O
Solvent: CH2Cl2
Time: 1 hour
Ambient temperature
Yield 70%
Ref. 1 5572812; Journal; Firouzabadi, H.; Sardarian, A. R.; Moosavipour, H.; Afshari, G. M.; SYNTBF; Synthesis; EN; 4; 1986; 285-288;

Does anyone have any comments on these methods?

How could Mn2O7 be prepared? from MnO2?? KMnO4? keeping it OTC....
Zinc dichromate wouldnt be to hard to make, adjust the normal dichromate synth method to replace KOH with Zn(OH)2 and this would produce the desired reagent.

runlabrun - 5-1-2005 at 02:09

lapse of memory....
2KMnO4 + 2H+ --> 2K+ + Mn2O7 + H2O

So...
2KMnO4 + H2SO4 --> K2SO4 + Mn2O7 + H2O
How could you seperate Mn2O7 from the post rxn mixture?

-rlr

BromicAcid - 5-1-2005 at 02:38

There is some information on Mn2O7 avalible in another thread here that you should read prior to attempts to use it in synthesis. It is essentially a covalent compound and is therefore not too soluble in the aqueous phase, a layer of it usually separates out. I would hardly consider this OTC though, it requires carbon tet which I've had a hell of a time trying to track down.

As for making ZnCr2O7, zinc hydroxide will not work as a replacement for potassium hydroxide, the properties of these two hydroxides are distinctly different and this substitution would be next to impossible. Zinc dichromate appears to be very soluble in cold water and decomposes in hot so that represents some prepratory problems. Mixing CrO3 with Zn(OH)2 may work in the presence of water. CrO3 readily made if you could pull off the dichormate synthesis in the first place with KOH provided you have acess to H2SO4.

THF + TCCA

skippy - 5-1-2005 at 09:56

Here's something from the usenet:
"
Juenge and Beal, in Tetrahedron Letters No. 55 pp 5819-
5820 (1968) described a method of oxidizing ethers to esters in one step.
They used trichloroisocyanuric acid (which (conveniently enough) is in some
swimming-pool chlorination chemicals) in a 3 molar excess of water to
oxidize a variety of ethers to the corresponding ester. Conditions were
3 degrees C, and a reaction time of 3 - 20 hours, depending on the specific
ether
"

I've read the paper mentioned and, sadly, it is no more in depth than the above quotation :(
Anybody have any thoughts on this method? The ethers oxidized in the paper were all non-cyclic, but maybe its worth a shot.

PS I found PVC glue at the home depot that contains only THF and polyester resin. Should be easy to separate. The price was around 10$ cdn for a little less than half a litre.

[Edited on 5-1-2005 by skippy]

[Edited on 5-1-2005 by skippy]

[Edited on 5-1-2005 by skippy]

runlabrun - 5-1-2005 at 14:17

Bromic:
I was thinking CCl4 wouldnt be to hard from the information in your thread on it. Calcium cyanamide from urea and calcium carbonate heating, and reaction with HCl:
CaCN2 + HCl(g) ---> CaCl2 + CCl4(g) + 2NH3
The OTC dichromate synthesis is easy to complete, i have used it many times, although not a very cost effective route... And access to conc H2SO4 is no problem whatsoever...
And thanks im reading the Mn2O7 thread now, i didnt seem to find it when i was searching here..... whoops!

Skippy:
Interesting method, anyone heard of it?
Worth a shot yeah but THF resources are thin unless you work in the lab, someone with easier access might want to try it...

-rlr

Mephisto - 5-1-2005 at 16:02

Skippy's method works also with the 1,4-butandiole, but I doubt that it is easier to get for you.

1,4-butandiole + TCCA to find in:
Synth. Commun., 25(5), 719-724 (1995) - Convenient Synthesis of Lactones by the Reaction of Diols with N-Haloamides
Syn. Commun., 33(12), 2003-2009 (2003) - The Oxidation of Primary Alcohols to Methyl Esters and Diols to Lactones Using Trichloroisocyanuric Acid

runlabrun - 5-1-2005 at 17:39

Ok, ive read the Mn2O7 thread and looked around a bit in my resources and texts.... i think i understand... i think.... :)

If conc sulphuric acid is added to a sep funnel and KMnO4 is added Mn2O7 will be formed as per the reaction mentioned above... would this be soluble in sulphuric acid or layer? The layer would be on the bottom right? 2.4g/mL for Mn2O7 compared to 1.8g/mL for H2SO4.
You could then add the Mn2O7 quickly to a distillation flask containing THF in CCl4/Acetone sitting in a dry ice bath (sublimating CO2 is -57oC the temp differece wouldnt matter would it?).

1- Does Mn2O7 detonate with CCl4?
2- In this proceedure were mixing acetone and Mn2O7.... boom? or does the CCl4 prevent this from happening?
3- CCl4 and Mn2O7 at these temps would be solids... would the rapid freezing of Mn2O7 from the room temp reaction cause it to detonate? or i could slowly cool the Mn2O7 and CCl4 in the flask in the dry ice bath and add the Acetone/THF mixture....

The details provide no times for the reaction, also no amounts are noted.... would it be:
2THF + Mn2O7 --> 2GBL + 2MnO2 + H2O ??
How much of each solvent?
So once this is complete the MnO2 can be filtered out of the mixture and GBL distilled to relative purity.

Sound reasonable? or a pipe dream?
-rlr

[Edited on 6-1-2005 by runlabrun]

cyclonite4 - 6-1-2005 at 04:27

skippy, there is some info in the THF thread about extracttion of THF from PVC pipe glue: http://www.sciencemadness.org/talk/viewthread.php?tid=2437

runlabrun - 6-1-2005 at 12:31

Thanks for the link.
I had already read the thread for the pvc solvent information however what i did just notice was the orgsyn link to 2-Furaldehyde.....

Could we go 2-Furaldehyde --> GBL if so by what reagent? the furfural would have to be saturated to get rid of the double bonded carbons and the aldehyde oxidised to form the lactone.

-rlr

skippy - 9-1-2005 at 10:41

Ok, the glue I was talking about earlier is called Oatey X-15 Bonding adhesive "for PVC sheeting". "CONTAINS: tetrahydrofuran, Polyester resin."
It comes in a blue pint can, and I found it at the home depot. I'll let you know how the separation goes

THF distillation

skippy - 26-1-2005 at 17:29

I washed the glue several times with about an equal volumes of water. The glue turned into a ball of putty floating in the water with the washing. The putty was discarded and the wash water distilled. The distillate's SG was measured to be around 0.88 or so, which matches with THF's density. I stopped the distillation early and spilled some condensate and still managed to get 200 mL of THF from half litre of glue. :)

[Edited on 27-1-2005 by skippy]

mick - 28-1-2005 at 13:16

I used to separate organolithium stuff from the THF organics by adding water to quench the reaction and then saturating with salt. Extract and then dry with anhydrous MgSO4. The THF was distilled off under vacum. I never noticed condensation and the samples seemed to be dry.
mick

skippy - 28-1-2005 at 15:30

Well, I've done two micro reactions and have learned that the TCCA should be added gradually. First test tube reaction: I added all of the TCCA in one minute, the TCCA had been powdered. Within less than a minute yellow chlorine gas was pouring out of the test tube and the THF and water were boiling out! I thought the problem was the powdering so I tried again with a single mini puck of TCCA and the reaction didn't start heating up until after 40 minutes or so. This time the reaction didn't boil over and only got up to around 40 C, so in the morning I took the post reaction mixture and evaporated it in a dish and was left with a rubber tasting goop. Interesting

newboy - 18-12-2006 at 06:13

Ive got 20 ltrs of THF and im in the process of getting potassium bromide can someone help me with the next steps

SecretSquirrel - 18-12-2006 at 11:46

You need potassium bromate for oxidation, but massive production of drug precursors is against forum rules, so I think you're on your own.

spanner - 20-12-2006 at 13:53

I ran across this information hope it help someone but i have a couple of questions related to it if anyone can help me

http://www.designer-drugs.com/pte/12.162.180.114/dcd/pdf/thf...

As i understand the process, the Ether is oxidized to the coressponding Ester Using NaBrO3 and NaHSO3, but it is not the actual reactants that directly oxidize the Ether, A redox is set up between the NaBrO3 and NaHSO3 were one of the resulting products are Br anions which do the actual oxidization of the ether.A chain reaction of BrO3 and Br is set up and as the reaction progresses it gets faster and self sustaining till the reaction completes itself.

My question is could the oxidization take place simply by using NaBr or would it need something else to push the reaction forward.

Nicodem - 20-12-2006 at 14:12

Quote:
Originally posted by spanner
As i understand the process, the Ether is oxidized to the coressponding Ester Using NaBrO3 and NaHSO3

According to the general procedure of the paper you linked the ethers get oxidized by HBrO3 formed in situ from KBrO3 and KHSO4 as acid. The reaction requires an induction period due to autocatalysis (since one of the products, Br2, is also a catalyst for the oxidation). NaHSO3 has no role whatsoever in the reaction as described in the general procedure (except for its use in the postreaction quenching of the oxidative species, but that is after the reaction itself).
Quote:
My question is could the oxidization take place simply by using NaBr or would it need something else to push the reaction forward.

NaBr can not oxidize anything since the bromide anions are already at the lowest chemically allowable oxidation state. The bromide anion can utmost act as a reducent but never as oxidant.

[Edited on 20-12-2006 by Nicodem]

spanner - 20-12-2006 at 14:20

thanks nicodem that help to clear things up a bit

spanner - 21-12-2006 at 01:08

Can the Br be recovered after the reaction is completed

Nicodem - 22-12-2006 at 00:54

Recovered as what? You can not recover it as KBrO3 because it gets used during the reaction.
Recovering it as KBr would be a bit too much work for such a cheap chemical. Similarly oxidizing the left over solid from the distillation would give Br2 which is a bit too hazardous to isolate for being worth the effort. In any case, if you would want to recycle it to KBrO3, then it would be a bit stupid to use KBrO3 for the oxidation in the first place, since you would need electrolysis for the reoxidation of KBr. Therefore it would be much simpler and cleaner to form gamma-butyrolactone from tetrahydrofuran/H2O in the presence of a small amount of Br2 or HBr by electrolysis instead.

kmno4 - 12-4-2008 at 06:27

Pffff....
It is so dusty in here :)
Using NaBrO3 in oxidation of ethers is interesting in itself.
I am wondering about much cheaper K/NaClO3 as replacement for bromate.
BTW>
According to the first given reference, NaHSO3 can be a real reagent for such reaction.
Bull.Chem.Soc.Jpn.,1997, 2561 - 2566 in attachment
Oxidation of Diols and Ethers by NaBrO3/NaHSO3 Reagent

Attachment: asd.pdf (221kB)
This file has been downloaded 4614 times

Nicodem - 14-4-2008 at 07:32

That paper is quite interesting as it describes a simple method for synthesis of bromohydrines without using N-bromosuccinimide or other such reagents. Surely an interesting method for BrO<sub>2</sub><sup>-</sup> in situ formation (bromites are rarely commercially available if at all).
However, for the oxidation of ethers the reaction stoichiometry is not as good as in the earlier mentioned Tetrahedron, 56 (2000) 1905–1910, and thus less ester is obtained per mol of KBrO3. On the other hand the reaction is probably much smother to upscale and should be less prone to runaways when compared to the Tetrahedron one which relays on an induction period (due to HBrO3 itself not being as efficient in oxidizing the alpha-position of ethers as Br2, HOBr and similar later forming species are).
Probably you would encounter the same problem using chlorates. They themselves are probably much less reactive toward ethers when compared to Cl(I) and Cl(III) species, so an induction period might be necessary to form enough reduced Cl species (or alternatively you might add a drop of HCl at the beginning to start the autocatalytic cascade right away). However, alpha oxidation of ethers with Cl2, hypochlorites, TCCA and the like are generally much messier than using comparable bromine based oxidants. All kind of nasty compounds can form and if your product is a liquid you might find it hard to purify it even with a good distillation column.

Methyl.Magic - 14-4-2008 at 08:36

Don't forget that lactones are easily synthetised by deshydrogenating diols.

Have anyone ever tried to hydrate cyclic ethers to diols ??

Nicodem - 14-4-2008 at 08:53

I don't know if it is possible, but the last time it was discussed I could provide no one step method to transform cyclic ethers like THF into terminal diols:
http://sciencemadness.org/talk/viewthread.php?tid=7765&p...
(though it's not like I searched the literature thoroughly enough to be sure of not having missed something simple)

sonogashira - 23-4-2008 at 06:09

I don't know if this has been mentioned before but since I came across the information I thought I would share:

THF oxidation to GBL with Ca(OCl)2 (63%) - Chem. Eur. J. 2007 13, 167-177 (Experimental on page 173 for preparing hexadeuterio analog of GBL)

and same reagents with 68% yield - Tet. Lett. Vol. 23, No. 1 pp. 35-38 (1982)

Would attach the papers but i'm having trouble with that for some reason.
I haven't tried either of the procedures since GBL is legal in my country. Best of luck!

kmno4 - 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.

kmno4 - 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.

stateofhack - 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]

Nicodem - 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.

Claret

ScienceSquirrel - 18-6-2008 at 09:29

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...

stateofhack - 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.

woelen - 18-6-2008 at 09:47

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

stateofhack - 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

chemoleo - 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?

MagicJigPipe - 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.

Panache - 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

The_Natural - 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.

ScienceSquirrel - 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.

The_Natural - 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.

ScienceSquirrel - 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

ScienceSquirrel - 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

The_Natural - 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.

chemoleo - 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.

ScienceSquirrel - 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...

The_Natural - 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.

MagicJigPipe - 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?

ScienceSquirrel - 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.

The_Natural - 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.

halogenstruck - 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.

Methyl.Magic - 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.

ergoamide - 14-12-2009 at 03:57

That's a bit of a general statement. Sure some are just cooks but some do have a real interest in chemistry. Besides most of them will already know the method you've mentioned. It's not like it hard to find a google search will get it for you.

Muzz1969 - 22-12-2014 at 04:47

Quote: Originally posted by Nicodem  
Recovered as what? You can not recover it as KBrO3 because it gets used during the reaction.
Recovering it as KBr would be a bit too much work for such a cheap chemical. Similarly oxidizing the left over solid from the distillation would give Br2 which is a bit too hazardous to isolate for being worth the effort. In any case, if you would want to recycle it to KBrO3, then it would be a bit stupid to use KBrO3 for the oxidation in the first place, since you would need electrolysis for the reoxidation of KBr. Therefore it would be much simpler and cleaner to form gamma-butyrolactone from tetrahydrofuran/H2O in the presence of a small amount of Br2 or HBr by electrolysis instead.



Hi Nicodem,
I am really interested in your electrolysis comment and was wondering can you explain the reaction mechanism , I'm just starting to learn about electrolysis reaction and I understand basic reactions but getting into organic reactions are a little above my head. Not interested in actually trying it just understanding how it works better.