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Author: Subject: benzyl alcohol oxidation
Magpie
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[*] posted on 7-5-2008 at 15:12


Perhaps a formula for Bz, i.e., "benzoyl," would help as an explanation:

Bz = C6H5-(C=O)-

e.g., benzoyl chloride, C6H5-(C=O)-Cl

BzH then is C6H5-(C=O)-H, or more commonly, C6H5-CHO

[Edited on 7-5-2008 by Magpie]
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[*] posted on 19-5-2008 at 19:22


Well, based upon the extremely strong smell of BnH 2 hours since the start of the reaction, I'd say the reaction kmno4 posted does indeed work.

It also scales up well, self shielding due formed N2O, and if one uses a bubbler it even lets you know when its done.



[Edited on 19-5-2008 by evil_lurker]




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[*] posted on 20-5-2008 at 09:05


Nice setup! I'm jealous :)



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[*] posted on 10-6-2008 at 11:14


I wanted to ask kmno4 a question about his comment of the HCl being the "dense lower layer". When I tried this a while back, that did not seem to be the case.

Wouldn't the BnCl sink to the bottom as long as the HCl (acid) was below ~20% concentration or so?

Quick check of Wikipedia:
38% HCl solution density: 1.18g/mL
BnCl density: 1.1g/mL

This is apparently at room temperature.

I found a list of HCl densities and at 20*C, 20% HCl's density (1.09g/mL) is less than BnCl's.

So, in theory, as long as your HCl was less concentrated than >%20 it should float to the top.

Sources:
http://en.wikipedia.org/wiki/Hydrochloric_acid
http://en.wikipedia.org/wiki/Benzyl_chloride
http://www.solvaychlorinatedinorganics.com/docroot/chlo_inor...




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[*] posted on 10-6-2008 at 11:40


Seems to me I recall my professor stating she never believes a synthesis when it suggests which layer is aqueous. ;) Always test.

Tim




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Fleaker
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[*] posted on 10-6-2008 at 12:48


MJP, that is what I said long ago from when I tried the reaction of benzyl alcohol to benzyl chloride. kmno4 didn't agree with it since his experience was that the lower layer was the acid. I never saw that. I lost all interest in argument and discussion when he put up his test results from a test in which he did not even reflux the reactants!! That explained it all to my satisfaction! :-) My BzCl sunk to the bottom because I did not use a gross excess of conc. HCl, and also consider well that the conc. of HCl decreased as it reacted with the alcohol. In my case, the benzyl chloride was always at the bottom. As kmno4 found out, that may not always be the case, depending on the amount of HCl you use, and whether you actually heat it like you are supposed to do! :D



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[*] posted on 10-6-2008 at 23:00


Quote:
Originally posted by Fleaker
I don't know what the Keq is for the reaction of PhCH2OH and HCl but I do know that the reaction is practically quantitative in its yield of benzyl chloride. It is a quick and efficient reaction. The acid is the solvent for the alcohol and is in great excess. Everything should be stirred well for several hours (that is probably too long). The dense lower layer is collected and separated with a funnel to remove any of the acid solution (...)

Where is your refluxing ? Do not play a fool, man.
Besides, I seriously doubt that refluxing HCl(aq) with benzyl alcohol is good solution for making benzyl chloride. This procedure is not recomended by some unavailable reference (it is unavailable because of MagicJigPipe and his banning at solo's wet dreams - info by solo :P). Benzyl alc. is sensitive to heating with acids ( info from available online wiley's encyclopiedia): heating same alcohol with acid can be dangerous, because of its spontaneous polycondensation. I think that in case of heating with conc. HCl such a condensation take place, at least partially. And, as I have already written, refluxing conc. HCl does not belong to pleasure things.....





[Edited on 11-6-2008 by kmno4]
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[*] posted on 11-6-2008 at 03:12


I have not seen procedures where reflux was used. Generally it is done at room temperature or with just slight warming (it can even be done directly in a separatory funnel). It is counterintuitive to use reflux for reactions that are rapid enough already at room temperature since the increase in temperature often promotes side reactions.
Refluxing conc. HCl results in expelling excess HCl gas untill negative azeotrope forms which is 20% HCl(aq). So it is expected that in such case the lower layer is benzyl chloride.
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[*] posted on 4-9-2009 at 12:10


The Oxidation of BnOH to BzH can be carried out by the reaction of the alcohol with gaseous NO2 at room temperature in a closed vessel. After some hours, the formed nitric acid and nitrous oxides are evacuated off and pure BzH is left. Procedure in ChemSusChem 2009, 83-88.

Has anyone tried this yet?
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[*] posted on 23-7-2015 at 05:26


Inspired by Magpie's original procedure, we've been working on refining the benzyl alcohol / persulfate oxidation.

It's a tricky one to get right but we can now boast an 85% yield. The trick is controlling the temperature very carefully, not too high so that it is out of control and over-oxidised, but also not going too slowly and therefore risking slow over-oxidation of product. You need to experiment and find the exact temperature (for your set up / solution) which is the 'initiation' point for the radical reaction, and steadily hold the reaction at this point by carefully adding the persulfate with vigorous stirring. Interestingly the occasional transient temperature rise due to too much addition doesn't hurt the yield too much. Above 72-75 degrees C however the yields are lowered due to oxidation to benzoic acid.

We've done a video showing the process and workup:

https://www.youtube.com/watch?v=YQWD4YIKBl4




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[*] posted on 23-7-2015 at 10:48


Very nice! You folks really took the time to investigate the nuances of this synthesis.

How did you conclude that this is a radical reaction? Can you write the equation for the reaction?

You achieved a significant improvement in yield. But likely you did not get the purity of my procedure. You used the bisulfite adduct. I used a steam distillation followed by vacuum fractional distillation. My product is clear and I obtained a boiling point indication of purity.

I also think that my procedure is somewhat safer than yours. By keeping the product hot enough there is an immediate and complete reaction when the drops of benzyl alcohol are added. There is no risk of a buildup of benzyl alcohol with consequent runaway. I recall very little char, and this is left in the pot during the distillations.

I would not have thought this reaction would be possible thinking instead that the product would be benzoic acid. We have to thank NERV and Fleaker for first mentioning this. I don't know where they got their inspiration. It's too bad they don't post anymore. Fleaker is a highly experienced/knowledgeable professional chemist in the metals area - NERV was a ChE student.

Is Tess married?

[Edited on 23-7-2015 by Magpie]

[Edited on 23-7-2015 by Magpie]




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[*] posted on 23-7-2015 at 18:26


Thanks Magpie - all good points. Yes, we only discovered this variation through curiosity and really playing around, and it took a lot of tries.

For anyone starting out I'd have to recommend starting with your process (and to follow it to the letter exactly), and then if you want to be more adventurous to give this a go.

The first time we tried this we followed your exact process and got a yield of around 50%. This is repeatable if you follow the exact process, but we noticed pretty quickly that even a small increase in temperature and over-oxidation starts to occur, and the most important factor seems to simply be time ; more time in the pot = more product oxidation. It's a delicate balance between a few different variables.

The intent is to vacuum distill the product once our pump is repaired (currently works for filtering but won't hack it for distillation) but you can see from the quantity of adduct that the yield is real. We didn't believe it first time and spent hours re-converting and washing the adduct, and trying to boil off DCM which we thought must be in there...

On the safety front we did do a small worst-case 'Trinity' test which was to saturate a 50:50 methanol-water mixture with persulfate and benzyl alcohol, and then slowly heat the homogenous mixture using a hot water bath. Nothing happens for a while, and then the reaction goes runaway in the space of about 20 seconds (starting at around 60 degrees). It was 'very strong' but not violent, and we concluded from this that so long as you've never got more than about 25ml of persulfate as excess unreacted in the flask you're only risking a slight reflux for a few seconds given the volume of liquid in there before it dies back down. Right at the start is the risky point, and the additional water and the methanol 'temperature buffer' help - the temperature in event of a small runaway reaches about 80C rather than the 110-120C if you don't have it present.

Agree it would be great to find out the exact mechanism of how this works; the way that the reaction goes feels like it's radical based, and persulfate is a known initiator, but of course this isn't evidence. Another clue is the fact that it works at neutral pH (it's acid conditions that favour the persulfate anion acting as the oxidising agent). Unfortunately we have no classical radical initiator compounds lying around we could use to see if this kick-starts the reaction, which would give some evidence.

There is some good research out there on the use of persulfates to degrade organic materials as part of environmental processing, and we're looking at this to give us ideas as to where we take this.

We'll try doing this in slightly alkali conditions (sodium carbonate to help prevent Cannizzaro) as this is known to assist radical formation in persulfate solutions, but typically this is at a pH>10 and our guess is that the pH and possibly the more reactive radicals formed will rapidly destroy the benzaldehyde formed. It might offer evidence of a radical mechanism though if the reaction proceeds spontaneously or at a lower temperature... Chelated iron salts are another unconventional initiator option we could try here.

We think this is worth playing with because it's feasible in theory that a radical oxidation mechanism with the right conditions would be quite powerful and might even open up an aqueous conditions pathway from toluene. Long shot, but worth having a go...

Tess isn't married but her jurisdiction sadly does not afford Silicon based life-forms this opportunity!




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[*] posted on 24-7-2015 at 09:33


Quote: Originally posted by Nicodem  
You also forgot the context of the article – the authors were looking for an economical use of Bn2O which is a worthless side product in the industrial preparation of BnOH.

I was wondering,if the ether was worthless,it should be quite easily available, cheap as well as less dangerous compared to dimethyl ether.Could it be used to run grignards ?the solubility of other organic compounds in Bn2O would have to checked first.I will do that tommorrow.;)

[Edited on 24-7-2015 by CuReUS]
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[*] posted on 25-7-2015 at 05:10


I'm supprised nobody mentioned chromyl chloride + toluene --> benzaldehyde. Chromyl chloride oxidises toluene to benzaldehyde but not to benzoic acid.

chloride + chromate + sulphuric acid -distill-> chromyl chloride
chromyl chloride + toluene ---> benzaldehyde

Although the second reaction is quite tricky to do.
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[*] posted on 25-7-2015 at 05:54


Quote: Originally posted by MeshPL  
I'm supprised nobody mentioned chromyl chloride + toluene --> benzaldehyde.

Quote:
Then 40 g toluene in 2X its volume of CHCl3 was added slowly with stirring to 140 g CrO2Cl2 in 200 ml CHCl3 on a cold water bath. After standing overnight, H2SO3 was added with stirring, the whole was steam distilled, CHCl3 evaporated, and a 44% yield of benzaldehyde was isolated via bisulfite. Their description of this Etard looks different than others that I have seen in the journals.

from http://www.sciencemadness.org/talk/viewthread.php?tid=2223&a...

[Edited on 25-7-2015 by CuReUS]
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[*] posted on 13-8-2017 at 05:37


I read all the posts and I'm really surprised with the fact that nobody, except Leu, talked about the OTC Hypochlorite oxidation of benzyl alcohol.

Oxidation of benzyl alcohol with calcium hypoclorite (HTH pool) and alumina (Al2O3) affords 99% benzaldehyde, just stirring the reagents at room temperature for 4 hours, as said in the study i bring to the community as an attachment.

Also, if you are in a hurry to get your benzaldehyde, you can react the same reagents (benzyl alcohol, calcium hypochlorite and moist alumina) in a microwave oven for just 01 (one) minute, to afford 98% benzaldehyde, as said in the another paper i'm attaching.

Sorry guys, mainly Magpie, whose job done with persulfate oxidation of benzyl alcohol is respectable, but I think Hypoclorite oxidation of benzyl alcohol is more OTC and high yielding than all the methods I've seen at this whole thread and prepublication section.

Here's the papers below:

Attachment: alcohol oxidation by CaOCL to aldehydes.pdf (161kB)
This file has been downloaded 552 times

Attachment: Microwave Assisted Selective Oxidation of Benzylic Alcohols with Calcium Hypochlorite under Solvent-Free Conditions.pdf (62kB)
This file has been downloaded 456 times
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[*] posted on 13-8-2017 at 06:06


That's kind of interesting. I have never tried making benzaldehyde, but I suspect the process is rather finicky, especially with an oxidizer as strong as calcium hypochlorite. I'm not sure what the role of alumina is in those procedures....



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[*] posted on 13-8-2017 at 06:21


Quote: Originally posted by Chemi Pharma  
Sorry guys, mainly Magpie, whose job done with persulfate oxidation of benzyl alcohol is respectable, but I think Hypoclorite oxidation of benzyl alcohol is more OTC and high yielding than all the methods I've seen at this whole thread and prepublication section.

Look at the scale of those reactions, then imagine scaling up to... oh, let's say one liter of benzyl alcohol. And to make things interesting, let's try the microwave-assisted variant.

Hopefully you haven't found this out the hard way yet, but the reaction you're proposing is highly prone to thermal runaway at anything approaching useful scales. Truth is, oxidizing benzylic alcohols to benzaldehydes is perhaps the easiest oxidation in organic chemistry, and many a PhD dissertation has been written on new methods for doing it. The hard part is doing it at a useful scale with OTC reagents and an easy workup.

Personally, I've found KMnO4 to work really well, since permanganate will rapidly oxidize a lot of the alcohol to benzaldehyde initially, and MnO2 will selectively oxidize the remaining benzyl alcohol. Even though KMnO4 is able to oxidize benzyl alcohol all the way to the acid, that's only a minor side-reaction if you work out the stoichiometry correctly. By the time there's an appreciable amount of benzaldehyde to react into benzoic acid, the permanganate has mostly decomposed into weaker oxidizers.




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[*] posted on 13-8-2017 at 08:13


Quote: Originally posted by JJay  
That's kind of interesting. I have never tried making benzaldehyde, but I suspect the process is rather finicky, especially with an oxidizer as strong as calcium hypochlorite. I'm not sure what the role of alumina is in those procedures....


JJay, I don't think that hypochlorite is so strong oxidiser as persulfate, KMNO4 or nitric acid. May be the Alumina role in this reactions is act as a cataliser, avoiding an eventual runaway like Melgar said, giving a milder reaction instead with hypochlorite alone.


Quote: Originally posted by Melgar  
Look at the scale of those reactions, then imagine scaling up to... oh, let's say one liter of benzyl alcohol. And to make things interesting, let's try the microwave-assisted variant.

Hopefully you haven't found this out the hard way yet, but the reaction you're proposing is highly prone to thermal runaway at anything approaching useful scales. Truth is, oxidizing benzylic alcohols to benzaldehydes is perhaps the easiest oxidation in organic chemistry, and many a PhD dissertation has been written on new methods for doing it. The hard part is doing it at a useful scale with OTC reagents and an easy workup.


Melgar, please, give it a try! I've never read at nowhere, this reaction isn't scalable. Don't blame the authors if they only did a small batch run and ommited about the scalability of the reaction.

I think some of us could run a 500 ml benzyl alcohol batch and tell about the results. I really don't expect a runaway! As i just have written to JJay, I think the Alumina presence makes the reaction milder than with hypochlorite alone.

Of course I agree with you that careful is necessary with microwave assisted reactions, mainly while scaling a reaction involving flamable reagents like benzyl alcohol. It's not commercially feasible, but it may be produce enough to a home made Lab necessities.

Quote: Originally posted by Melgar  
Personally, I've found KMnO4 to work really well, since permanganate will rapidly oxidize a lot of the alcohol to benzaldehyde initially, and MnO2 will selectively oxidize the remaining benzyl alcohol...


Ok, and what about the yields of benzaldehyde in this KMNO4 oxidation? I have my doubts if it not results in a mixture with near equal quantities of benzaldehyde and benzoic acid due the powerfull oxidating properties of permanganate.



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[*] posted on 13-8-2017 at 09:23


Hmm... I had thought hypochlorite was a stronger oxidizer than almost anything except hydrogen peroxide... this table doesn't really say: http://chemed.chem.purdue.edu/genchem/topicreview/bp/ch19/ox...





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[*] posted on 13-8-2017 at 09:58


Take a look here JJay:

http://hyperphysics.phy-astr.gsu.edu/hbase/Tables/electpot.h...

As you can see, hypochlorite anion has an eletrochemical potencial of + 0,90V, while HNO3 has + 0,96V, Chlorine + 1,36V, KmnO4 + 1,49V, H2O2 + 1,78V and persulfate + 2,01V.

So, hypochlorite has the minor oxidant power between all this reagents above.

[Edited on 13-8-2017 by Chemi Pharma]
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[*] posted on 13-8-2017 at 10:14


Isn't there more to oxidizer strength than electrode potential, though? Dichromate has lower electronegativity than hypochlorite, but chromium +3 is oxidized to chromium +6 by hypochlorite....



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[*] posted on 13-8-2017 at 10:36


It's cause the oxidant agent able to oxidate Cr+3 to Cr+6 is the free chlorine produced by the dissociation of the hypochlorite anion at low PH (below 7), nor the hypochlorite ion.

As you can see at the table, free Chlorine has a redox potencial of + 1,36V, while Cr+6 has 1,33V. This little difference may explain why chromium salts are only slowly oxidated to Cr+6 during chlorination of drinking water.

I found a theme at PubMed telling about the fact. Take a look:

Oxidation of Cr(III) to Cr(VI) during chlorination of drinking water

https://www.ncbi.nlm.nih.gov/pubmed/22487808

No way the oxi-redox table were wrong JJay, otherwise it would be a breakdown of all the expected phisicochemical rules.
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[*] posted on 13-8-2017 at 11:18


I've oxidized Cr+3 to Cr+6 at high pH, so I'm not really sure what you're getting at. I don't think anyone said the table is wrong.



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[*] posted on 13-8-2017 at 12:04


Quote: Originally posted by JJay  
I've oxidized Cr+3 to Cr+6 at high pH, so I'm not really sure what you're getting at. I don't think anyone said the table is wrong.


Ok JJay, I apologize for the misunderstanding. Let me try to explain in the right way:

1- It's not possible a reaction with eletrochemical potencial lower than other displace the equilibrium in favor of the first. To do this we need to force some variables like pression, temperature, eletrochemical charge input, like electrolisys, etc.

It's the same principle behind the reaction of metals more eletropositive than hydrogen reacting with acids to give hydrogen gas (displacement), while metals like silver and copper doesn't.

The eletrochemical potencial of a reaction can tell us about the strenght of the oxidazing or reducting power of a reagent, cause the two things it's intimally related.

What's facilitate an oxi-redox reaction? the hability of the reactant to receive electrons and the reagent to give them. Then, to know if some reaction will be spontaneous or not, just look at the electrochemical potencial of the equation, based on the table.

2 - Therefore, when we oxidate Cr+3 to Cr+6 with hypochlorite, at really, the residual molecular free chlorine in solution is who do the job, not the hypochlorite anion, cause ClO- to Cl is more eletropositive than Cr+3 to Cr+6. That is impossible.

Off course free molecular chlorine exists dissolved in a hypochlorite solution, even with PH 14. I just said that if you diminish the PH, it favours the formation of hypochlorous acid (HCLO), releasing more free chlorine in solution, increasing the oxidating power to transform Cr+3 in Cr+6.

Forgive me if you feel offended by the way I expressed my ideas. It was not my intention.


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