Sciencemadness Discussion Board

Benzene synthesis

 Pages:  1    3  

S.C. Wack - 25-7-2009 at 10:33

Looked up reference 2 in the paper that I linked to - in Chem. Lett. 33, 282 (2004), the most ideal conditions for benzene in that system were: PET powder was heated with Ca(OH)2 at 700C at a 10:1 Ca(OH)2 : PET ratio, giving a distillate that was 45.5% of the weight of the PET, and benzene accounted for 36% of the weight of the PET.

They analyzed the products closely and quantitatively under the various parameters used, and quite a number of things were present. Under the ideal conditions, the highest % contaminants of the distillate were, again in wt.% of PET: 3.5% biphenyl, 2.3% toluene, 1.36% styrene, and 0.6% acetophenone. Using any less, and especially much less Ca(OH)2, dramatically decreased the selectivity for benzene.

[Edited on 25-7-2009 by S.C. Wack]

JohnWW - 25-7-2009 at 20:38

But, on an industrial scale, how would such a process for obtaining benzene, based on hydrolysis of PET (probably from recycled plastic drink bottles, most of which BTW are shipped to China) and decomposition of the monomeric products thereof, compare for unit co$t and efficiency with either fractionation of benzene from petroleum, or direct synthesis of benzene by polymerization of acetylene under heat and pressure? (The acetylene, C2H2, having been generated by hydrolysis of calcium carbide, CaC2, obtained by pyrolysis of lime with coke or anthracite in a retort). As a chemical engineer, I am aware of such considerations.

Zinc - 14-9-2009 at 07:14

Sorry if this is a little off topic, but is there any way toluene could be obtained by some pyrolysis procedure like benzene can?

Picric-A - 14-9-2009 at 09:02

Yes, heating 2-methyl-benzoic acid with NaOH will produce toluene

Eclectic - 8-11-2009 at 08:55

Does anyone know if benzenesulfonic acid would distill (190 C) from benzoic acid in hot sulfuric acid as a decarboxylation solvent? Perhaps with a little copper as catalyst?

manimal - 6-4-2010 at 15:12

Apparently, Org recently passed away from cancer. One wonders if his fondness for chemicals such as this played a part.

Jor - 6-4-2010 at 16:13

Org? Organikum? What kind of cancer?

Whoever it was, rest in peace!

manimal - 6-4-2010 at 17:48

Quote: Originally posted by Jor  
Whoever it was, rest in peace!


Seconded.

Organikum, yes; I'm not sure of the details, that's just what I read on another forum (WD).

hissingnoise - 7-4-2010 at 03:41

It's saddening to note that our clandustrial priest won't be posting again.
I hope he enjoyed his life.

S.C. Wack - 7-4-2010 at 11:31

IIRC he mentioned a terminal lung cancer diagnosis (he was a smoker) in passing in his deliberately last post there in a threat --- about 5 months before his last post here. Note how the last login here is much later. IDK if he said anything in whimsy here that no one noticed, if he hated the RIP Halfapint thread at the-hive, or what the deal is. Anyone that knows for sure isn't saying.

unionised - 25-4-2010 at 12:38

I just looked through this thread and didn't see a reference but has anyone considered the pyrolysis of polyvinyl acetate (or alcohol) as a synthetic pathway to benzene?

Pyrolysis of PVC works too, but I'd not be happy about the potential for making chlorinated aromatics.

Not sure how well this link will work, but it indicates that benzene is a major product (based on gc peak area which isn't a brilliant way to quantify stuff)
http://books.google.co.uk/books?id=AY__YWnIE_4C&pg=PA312...

There's a table of products on p326.

[Edited on 25-4-10 by unionised]

densest - 25-4-2010 at 13:21

Sort-of-on-topic: I heard about a company in Houston which is using large plasma reactors to reprocess waste such as PET, other plastic scraps, and about anything else which can be vaporized in their chambers, the output being fractionated with (I am guessing) MHD or other very large scale electric-magnetic-mass-??? separators.

I'm trying to track this down... google didn't have anything suggestive last time I looked. Any rumors anywhere?

not_important - 25-4-2010 at 20:57


http://en.wikipedia.org/wiki/Plasma_arc_waste_disposal

http://www.popsci.com/scitech/article/2007-03/prophet-garbag...


While aromatics are formed in high temperature mixes rich in carbon and hydrogen in roughly unity ratio, SFAIK the yield generally aren't great and you tend to get mixtures.






Sandmeyer - 26-4-2010 at 03:09

Quote: Originally posted by S.C. Wack  
IIRC he mentioned a terminal lung cancer diagnosis (he was a smoker) in passing in his deliberately last post there in a threat --- about 5 months before his last post here. Note how the last login here is much later. IDK if he said anything in whimsy here that no one noticed, if he hated the RIP Halfapint thread at the-hive, or what the deal is. Anyone that knows for sure isn't saying.


Is it true that he is dead? If it is true then it is a great loss to internet chemistry forums. I truly miss his posts and his sense of humour, to hear that he is no longer here is very saddening.

majortom - 8-5-2010 at 08:26

Quote: Originally posted by BASF  
Do you have more information on the decarboxylation of benzoic acid?
Decomposition temp. and similar things?

I think the decarboxylation of benzoic acid would be very attractive for the home chemist due to the widespread use of benzoic acid in the food industry, and therefor it should be easy to get and not that expensive.
I still have some 100-200g of benzoic acid from a drug store.

HLR


Wikipedia has an interesting article on benzene in soda, not very much mind you but it is still benzene. It was apparently caused by decarboxylation of benzoic acid.

http://en.wikipedia.org/wiki/Benzene_in_soft_drinks

I have heard of thermal decarboxylations on amino acids, I have heard that ketones act as catalysts and increase yields, It used Tryptophan or Lysine in xylene with a dash of MEK (2-butanone) I am not sure what yields are or if the ketone does anything at all. Although the wiki article says citric acid and some other things in citrus soda's catalyze the reaction of benzoic acid.

[Edited on 8-5-2010 by majortom]

[Edited on 8-5-2010 by majortom]

not_important - 8-5-2010 at 08:57

that has been brought up before, the amounts formed are microscopic.


majortom - 8-5-2010 at 10:47

Diffidently, only a few micrograms but if the reaction can take place, albeit slowly, in soda in a very cool temp(even if it is 100 F) than what would happen if we heated it up to 150 C in some high boiling point solvent?

[Edited on 9-5-2010 by majortom]

testimento - 20-10-2013 at 04:45

I am unfortunate to not to be able to find data which could indicate how xylene or benzene could be transferred to toluene. Does anyone have any data or sources for this process?

I found one possible route:

p-Toluic acid is first made from xylene with KMnO4, and then it is decarboxylated with 2 moles of NaOH, leaving toluene and sodium carbonate residues.

[Edited on 20-10-2013 by testimento]

amazingchemistry - 20-10-2013 at 07:29

It might be possible to form benzene via decarboxylation of benzoic acid. The main article cited in wikipedia, from the Journal of Agricultural and Food Chemistry, does not make a thorough effort to find the best conditions for benzene synthesis, as it's goal is to find out whether decarboxylation occurs in realistic beverage storage and consumption conditions. And indeed it does, albeit in microscopic amounts, as has already been pointed out. I'm curious if yield could be improved by heating/refluxing, by finding a better decarboxylating agent (ascorbic acid is used, citric acid is also mentioned but this produces even lower yields), or by finding a better transition metal catalyst (mainly iron and copper compounds are used). I've had this in mind for months now, but unfortunately, I can't afford the equipment or reagents to test this, or find a suitable place to carry out these experiments.

DJF90 - 20-10-2013 at 10:14

Quote: Originally posted by amazingchemistry  
It might be possible to form benzene via decarboxylation of benzoic acid.


Have you not bothered reading this thread?? Benzene was fairly easily obtained by this route. A friend of mine did this to make over 3 L of benzene.

amazingchemistry - 20-10-2013 at 20:54

By "fairly easily" do you mean by the high-temperature solid-solid distillation of benzoic acid and sodium hydroxide? The same distillation where you have to be very careful temperature-wise or your glassware melts on you (molten hot NaOH being death to glass)? Do you mean the same distillation that because of this, more often than not, requires all-metal setups (many of which are pictured in this and other threads)? Do you mean the same distillation that yields an unsightly red product that you have to redistill to get anything close to pure benzene? Or have we made progress in the couple of months I've been away from this forum? If you bothered to read the paper in question you'd know that this particular decarboxylation was achieved in an aqueous solution of sodium benzoate, maintained at pH 2 with a buffer, using ascorbic acid as a hydroxyl radical donor and a transition metal catalyst. There was some decarboxylation going on even at room temperature. No need for homemade metal setups, crazy temperatures, etc. Although the reported yield is very small at typical beverage storage and consumption conditions, the paper did make me think about whether efficient benzene synthesis can be achieved by this route, given optimized reaction conditions. This is a valid research question, but unfortunately I have neither the equipment nor the place to pursue it. Maybe others can?

[Edited on 21-10-2013 by amazingchemistry]

Lambda-Eyde - 20-10-2013 at 21:00

Decarboxylations can be done at lower temperatures with other reagents. Here is an example. Although I see no reason to abandon the current established route, as NaOH is dirt cheap and readily available and the procedure itself is piss easy for anyone who has a spare paint can.

amazingchemistry - 21-10-2013 at 09:51

I don't see a reason to abandon the current route either (after all, we know it works). However, as far as I know, no experiments have been done to explore whether the above route or other routes would be feasible, or perhaps even better than the current one. I'm interested in the route described above because both ascorbic acid and sodium benzoate are cheap and readily available, and because the procedure described in the paper amounts to little more than 'mix and heat.' If that doesn't work, maybe we can learn something from it. We're chemists after all, and most of the stuff we do doesn't work :D

DJF90 - 21-10-2013 at 12:44

Quote: Originally posted by amazingchemistry  
...procedure described in the paper amounts to little more than 'mix and heat.'


In such simple terms, thats exactly what the current method is.

testimento - 6-2-2014 at 15:25

What might be the easiest way to oxidize styrene, toluene or xylene into benzoic acid?

sargent1015 - 6-2-2014 at 20:26

Try this?
http://www.youtube.com/watch?v=iGTq43-2V_I

testimento - 7-2-2014 at 09:45

KMnO4 isn't available for me. I could make it, but when one needs 3-6 times the amount of that for 1 part of toluene, that's just way off scale.

How about this?

Benzoic acid is produced commercially by partial oxidation of toluene with oxygen. The process is catalyzed by cobalt or manganese naphthenates. The process uses cheap raw materials, and proceeds in high yield.

Should it need pressure, temperature? Max pressure I can reach is maybe 40-50bars with refrig. compressor. Could normal Cu salts catalyst work, or what is the role of napthenate in here?

EDIT

This looks most promising:

https://www.google.com/patents/US3210416

Essentially one will need a steel reactor pot where toluene is inserted with cobalt or manganese catalyst and possibly a promoter, and then the vessel is pressurized to 15-20bar and heated to 175-200C. Probably it should be mixed by jerking the pot around. This should be quite possible to produce with common fridge pump with one-way valve and copper pipes, a needle-relief valve and a pressure meter and thermoprobe on top.

But how does this differ from Dow Process of making phenol from toluene by oxidation (Toluene + 2 O2 = phenol etc.)?

[Edited on 7-2-2014 by testimento]

sargent1015 - 7-2-2014 at 10:13

Pressurizing vessels at home seems... Bad? Anyways, don't blow a hand off

Also, KMnO4:
2 pounds for $20

-Sarge

testimento - 7-2-2014 at 13:48

I have more experience in that area than in chemistry. :) Permanganates are banned in my country.

I read the patent now fully and it states much of the following:

Of total mass of toluene:
1 % manganese acetate
0.5 % ammonium bromide (i've got only sodium)
10-14bar pressure
140-180C temperature
Air bubbled inside

With high flow air it should form benzoic acid quite fast. In patent it is reported to produce several tons per hour with the specs mentioned. I think I could use ordinary compressor because the pressure range is from 4 bars to 20 bars. 8-10 bars should be pretty enough and since the reaction goes very fast, I can directly adjust the air flow with needle valve to compensate the consumed air(which should be mostly lack of oxygen).

Heating should be concerned, since the device is essentially a pressurized bomb which contains both oxidizer and fuel and it might go up pretty nice if it were to get fire. So electric heating should be preferred instead of gas burner, and the exhaust gas, which certainly contains vapors of toluene and other gunk, must be led into water trap or vented safely elsewhere or ignition may occur.

Attached picture for reactor.

The achieved benzoic acid could be turned into benzene or decarboxylated into phenol. What else could benzoic acid be used into?

potty.png - 14kB

[Edited on 8-2-2014 by testimento]

plante1999 - 17-4-2014 at 16:44

I just had the opportunity to get 80 ml of toluene by distilling the non-aqueous layer of an old varnish can. The first test I wanted to do is catalytic oxidation. It was just a first try using this principle.

I mixed 0.1 ml of bromine, a pinch of cobalt carbonate, and 30 ml of glacial acetic acid, I then added 20 ml of toluene and put the mixture in a 100 ml two neck flask. After the bromine addition the mixture was of a orange color. One neck was used to pass air using an aquarium pump in the mixture, and the other had two condenser inline setup for reflux. Heating was started and the mixture slowly turned dark red, up to a point it suddenly shifted to pink and air was passed in when the mixture suddenly changed color. Two condenser did not seem to be enough, as plenty of vapors could go all the way out of the apparatus. At some point, white fog formed on the glass which I presume is sublimed benzoic acid.

I runned it for an hour, and then I stopped. I still have yet to find a way to extract the benzoic acid without diluting the acetic acid, as it is quite hard to make/get in quantity.

blogfast25 - 26-4-2014 at 12:21

Here’s an exploratory experiment on benzene preparation from sodium benzoate and sodium hydroxide dry distillation.

30 g of sodium benzoate and 9 g of sodium hydroxide were ground into a fine powder with a granite mortar and pestle and loaded into the following ad hoc apparatus (here post reaction):



Because I’m not keen on using open flame with a highly flammable reaction product and rubber stoppers as sealant/connectors, I first tried heating on full on my electrical hot plate. No product came over.

Medium propane Bunsen heat immediately caused the reagent mix to char and orange distillate to form. I reckon this needs at least 400 C to get benzene to form. Also some white smoke that didn’t seem to want to condense formed. I kept heating somewhat intermittently until almost no white reagent mix could be seen. The large rubber stopper had by then also suffered badly and it was time to stop.

In all I got about 25 ml of orange product with the unmistakable smell of benzene.

Here’s a close up of the reactor, post reaction: mostly black gunge with some white material too;



For a larger scale test, a metal reactor or expendable glass reactor, will be needed. I think some of the above commenters are right about heat transfer: it appears difficult to heat the whole mass through.




[Edited on 26-4-2014 by blogfast25]

S.C. Wack - 26-4-2014 at 14:27

I'd specifically recommend Pyrex 4620/4680 dime a dozen flask auctions, and not exposing rubber stoppers to benzene, but this like everything else will just be ignored in 3 months if not now, so there's not much point to saying anything any more.

Charring = too hot, organic chemistry = low and slow.

blogfast25 - 27-4-2014 at 01:32

S.C.:

The glass I used was cheap and I think it will be recoverable.

Rubber stoppers: agreed but this wasn't an occasion where ground glass was available to me. I'll be using another type of sealant next time, possibly EPDM rubber with a much higher temperature/aromatics resistance.

Slow may be best but at an estimated 250 C (max setting on electrical plate) I got nothing whatsoever.

I think this preparation calls for a long, tubular type reactor geometry, placed on its long side and with gradual heating from one side to the other. This would also prevent thermally overloading the sealant.

[Edited on 27-4-2014 by blogfast25]

blogfast25 - 27-4-2014 at 13:12

Quote: Originally posted by Magpie  
Benzene making, 1st picture:

<img src="http://www.scimad.org/scipics2/Na_Benzoate_+_NaOH%20_2).JPG" width="800" />

<!-- bfesser_edit_tag -->[<a href="u2u.php?action=send&username=bfesser">bfesser</a>: reduced image size(s); hosted image(s) at scimad.org/scipics2/]

[Edited on 12.12.13 by bfesser]


Can anyone restore these photos?

plante1999 - 29-4-2014 at 08:21

I doupt so, did you re-distilled the crude benzene blogfast?

Magpie - 29-4-2014 at 09:06

Quote: Originally posted by blogfast25  
Quote: Originally posted by Magpie  
Benzene making, 1st picture:


<!-- bfesser_edit_tag -->[<a href="u2u.php?action=send&username=bfesser">bfesser</a>: reduced image size(s); hosted image(s) at scimad.org/scipics2/]

[Edited on 12.12.13 by bfesser]


Can anyone restore these photos?


I attempted to find these 2 photos on my hard drive but they are not there.

It appears that bfesser moved them from this thread to scipics for some reason. When I go to scipics I find that the domain is for sale. Does this mean that we have lost all the pictures on scipics forever?

[Edited on 29-4-2014 by Magpie]

blogfast25 - 29-4-2014 at 11:41

@plante: not yet, as the volume is so small and this was just about 'feeling the water'.

@magpie: where is bfesser when he needs a b*ll*cking? ;-)

Any chance of simply describing your set up?

I'm now thinking of a largish (30 mm OD), tubular reactor (boroglass, a large test tube in essence, I may have found just the thing on eBay), horizontally clamped with an EPDM rubber stopper and glass tube at the neck leading to the condenser. Charge to no more than half capacity, leaving 'head room' for the expanding charge. Gentle heating by moving a propane Bunsen flame along the tube, until crude benzene starts coming off.

I'm not keen on the many Utoob metal cans + 'some sealant' + copper tube: leakage of benzene could have very serious consequences with that set up...


[Edited on 29-4-2014 by blogfast25]

Organikum - 29-4-2014 at 12:09

Quote: Originally posted by blogfast25  

.....
I'm not keen on the many Utoob metal cans + 'some sealant' + copper tube: leakage of benzene could have very serious consequences with that set up...

[Edited on 29-4-2014 by blogfast25]


The point is that a standard thread of iron watertubing, you may know the conical one, fits snuggly into the opening of a solvent can where the plastic was removed. Some good amount of PTFE tape gives a sure seal. Just as always whe using PTFE tape -dont use force or youdestroy more then you seal.

I am now not going into details of your suggested nonsensical setup which is so prone to failure that it hurts only reaing it. Sorry, but its true.

If you want the last edge on saftey then fit a aspirator at the receiving vessel and run it just for a little suction. Cool receiver with ice, maybe a small condrensor with icewater on the flask before the aspirator connection.
This will speed things up and push yields a bit what will be lost by the aspirator again, but most important: no benzene nowhere never in the air.

If thats just a contest to suggest the most byzantine ways this might be implemented as the best ways are already known and just too simple, then I apologize.
I instead will consider to contribute in this sense myself.

regards
/ORG

Magpie - 29-4-2014 at 12:16

I think that a 30mm borosilicate glass tube 1/2 filled will work just fine. Keep in mind that the decarboxylation doesn't really go until the NaOH melts. Mp for NaOH is 318°C, so this will severely corrode the glass tube. But the tube can be reused, likely for several runs (5-10 maybe).

I don't make benzene anymore - I buy it. But if you just want to make it for education/fun I understand.

blogfast25 - 29-4-2014 at 12:32

@Magpie:

Hmm... the decarboxylation works also with Ca(OH)2 and that doesn't really melt.

The glass reactor I used was unaffected by the NaOH and after washing was 'as new'. This may be due to the emanating volatiles protecting the glass somewhat. Glass always has one advantage: you can see through it.

I'd gladly buy the benzene if I found a seller. No luck so far.

@Organikum:

"I am now not going into details of your suggested nonsensical setup which is so prone to failure that it hurts only reaing it. Sorry, but its true."

From your hyperbolic reaction I deduce you don't seem to understand what it is that I'm proposing. You owe it to me to explain what it is you believe that can go wrong with this 'nonsensical setup'.

Thanks for the tips on using cans, threads and Teflon.


[Edited on 29-4-2014 by blogfast25]

blogfast25 - 3-5-2014 at 10:25

I tested the idea of a tubular reactor today, using an old Nessler tube, OD about 30 mm and nominal capacity of 150 ml:



The charge was the same as in my first experiment: 30 g sodium benzoate and 9 g sodium hydroxide, ground up (about 50 ml), spread out over the length of the tube to allow easy heat transfer (only a thin layer of charge) and avoiding excessive expansion of the reagent mix.

I wanted to use a specially ordered black EPDM rubber stopper (better heat resistance) but it hadn’t arrived yet. So instead I used the usual (SBR?) orange stopper but this time wrapped carefully in Teflon gas tape. There’s several cm distance between the right hand side of the charge and the stopper, to try and keep the stopper as cool as possible.

I heated very carefully, constantly moving a medium propane Bunsen flame from one side of the charge to the other and back. But to get actual distillate (as well as white smoke also reported by Magpie) ultimately more heat was needed. I collected about 12 g of orange distillate, a yield based on benzene of 74 %. It doesn't appear to be possible to obtain benzene without charring the charge, at least not within reasonable time frames.

Here’s the tube after reaction:



Black patches, white patches and orange condensate on the top side.

I need to check the state of the rubber stopper to see if the Teflon wrap provided any protection.

The Nessler tube will be cleaned by soaking the inside in 30 % NaOH solution.


[Edited on 3-5-2014 by blogfast25]

HgDinis25 - 3-5-2014 at 10:45

http://web.anl.gov/PCS/acsfuel/preprint%20archive/Files/48_2...
The paper above states that, at 450º C, Sodium Benzoate may decompose to form a variety of producrs, including Triphenylmethane, wich is used to make a large number of dyes, the Triarylmethane dyes (http://en.wikipedia.org/wiki/Triarylmethane_dye).

This is a long shot, but is there any possibility that the Triphenylmethane is reacting with other products, thus making an orange dye? Anyway, most Triarylmethane dyes can be used as pH indicators because they change colour depending on the pH. Perhaps you could try to increase and lower the pH of your product to see if it changes colour?

blogfast25 - 3-5-2014 at 13:01

HgDinis25:

Quite a bit of work has gone into identifying the orange material, higher up in this thread. I can't remember the outcome, sorry. Thanks for the paper.

I'll take you up on the pH/colour idea but my hopes aren't high.

On a more mundane level, my Teflon wrapped stopper hadn't sustained any damage worth mentioning. That worked even better than hoped for.

[Edited on 3-5-2014 by blogfast25]

S.C. Wack - 3-5-2014 at 13:34

Quote: Originally posted by blogfast25  
It doesn't appear to be possible to obtain benzene without charring the charge, at least not within reasonable time frames.


Perhaps you're doing it wrong. For me there was a colorless distillate and no charring, until the heat was turned up when things tapered off, which did nothing worth doing and should not have been done.

blogfast25 - 4-5-2014 at 04:56

Quote: Originally posted by S.C. Wack  

Perhaps you're doing it wrong. For me there was a colorless distillate and no charring, until the heat was turned up when things tapered off, which did nothing worth doing and should not have been done.


I didn't see your contribution in this thread, point to it if you will?

The first drops came over clear and colourless but at an interminably slow rate. Even at higher rate it still took about 15 min for 12 ml. At that rate one is inclined to go full throttle at the expense of purity.

If I'm doing it wrong everyone else here, bar you (no sarcasm intended), seems to be too.

It would help greatly if someone could establish some rate-temperature relation for this reaction(s).


[Edited on 4-5-2014 by blogfast25]

blogfast25 - 4-5-2014 at 10:42

I also had a quick stab at the reduction of phenol with zinc powder. Using the same set up as above but slightly tilted so that the Nessler’s bottom was pointing downwards a bit. 10 g phenol and 8 g ultrafine Zn powder were then loaded into the Nessler tube.

I assume this reaction is supposed to go as:

Phenol + Zn === > Benzene + ZnO

Edit: hydrogen deleted as per deltaH (ta!)

I heated and got bubbles but no benzene. At higher heat the phenol started to distil over and froze in the condenser:




The zinc appeared also unaffected, still a grey dark powder.

It appears this reduction needs leading vapours of phenol over 400 C zinc granules or turnings, as mentioned in Wikipedia.

[Edited on 4-5-2014 by blogfast25]

blogfast25 - 4-5-2014 at 11:15

Quote: Originally posted by deltaH  

I don't think the hydrogen should be there, one hydrogen needs to go in the place of the -OH group


Oooopsie. You're right of course.

[Edited on 4-5-2014 by blogfast25]

plante1999 - 4-5-2014 at 14:22

You may want to try to pass phenol vapors over an heated tube containing zinc dust which would be heated using a bunzen.

blogfast25 - 5-5-2014 at 05:18

Quote: Originally posted by plante1999  
You may want to try to pass phenol vapors over an heated tube containing zinc dust which would be heated using a bunzen.


Sure, but it's a lot less practical as a lab preparation than a dry distillation: it requires quite a bit of engineering to do it safely, like condensing very hot benzene for instance. The latter would require a copper coil, I think. Any unreacted phenol would have to be dealt with too, in a separate cold trap.

Later on I'll have a quick test tube stab with aluminium powder.


[Edited on 5-5-2014 by blogfast25]

deltaH - 5-5-2014 at 06:39

blogfast, can you please ellaborate how you expect these reductions to proceed as I have a major concern:

Phenol is weakly acidic, so a reaction between phenol and an electropositive metal would most likely initially yield the metal phenate (presumable a solid) and produce hydrogen and heat, for example:

2Ar-OH(l) + Zn(s) => Zn(O-Ar)2(s) + H2(g) + heat

Idealy, one would reflux this to prevent phenol boiling off.

The problem is that further pyrolysis of the phenate is unlikely to yield benzene because you're now short the hydrogen having previously driven it off.

You might make diphenyl ether on pyrolysis, which would, nevertheless, be an interesting result :D

?? Zn(O-Ar)2(s) + heat => Ar-O-Ar(g) + ZnO(s) ??

Note that diphenyl ether is very high boiling (258°C), so it may condense near the pyrolysis heating zone.

Apparently it smells weakly of geraniums, though I doubt a pyrolytic product would smell very pleasant :o

****

I found this obscure and dated text about pyrolysing calcium phenates (not sure if phenolate or phenate is correct, the authors referred to it as phenate, so I'll stick to that). According to them, phenol yields an 'unusual' calcium phenate hydroxide 'half salt' upon reaction with calcium hydroxide which simply regenerates the phenol upon pyrolysis and forms CaO as co-product.

They also say that the pyrolysis of alkali phenates like sodium or potassium phenate yields only char and gas, but no volatile organic products (as per the finding of others, see references in article). That could be on account of the high basicity of those salts, so maybe some hope for zinc and aluminium to make something else?

See: https://web.anl.gov/PCS/acsfuel/preprint%20archive/Files/28_...

[Edited on 5-5-2014 by deltaH]

blogfast25 - 5-5-2014 at 10:13

DeltaH:

I really don't know and only go by a reference in wiki. Reduction of phenol with Zn has also been briefly discussed higher up. It's a question of Delta G, I think: the route leading to ZnO and not Zn phenolate is more exergenic. And w/o water phenol must be a poor proton donor: in molten (or gaseous) phenol no H<sup>+</sup> can probably exist in significant amounts.

Refluxing phenol would be awkward because of its MP (but not impossible).

I briefly contemplated phenolates too but there's that hydrogen problem...


[Edited on 5-5-2014 by blogfast25]

deltaH - 5-5-2014 at 13:34

Ah ok I understand now, I was thinking of it in too simplistic terms, you are right, at high temperatures in the gas phase, it might be a different ball game altogether.

However, that said, I can't shake the feeling that beside for the reductive abilities of the metal, its ability to transfer the hydrogen or more specifically, its kinetic ability to work with hydrogen is important. In that sense, I don't think zinc is particularly known to be catalytic with hydrogen nor form hydrides AFAIK, so may not be the best choice here (note, not saying it doesn't work ;) ).

In this regard, I think something like commercial titanium hydride powder might work well. My thinking here is that unlike titanium powder, which is probably too kinetically inert, the hydride is much more amorphous and active and so provides both hydrogenation ability and strong reducing ability in one package. I know that you don't theoretically need the extra hydrogen stoichiometrically, but I think for kinetic reasons, it would help a lot to have an excess for this reaction.

That said, I don't think benzene itself would hydrogenate (forming cyclohexane for example) as that is particularly tough and probably needs specialised hydrogenation catalysts like palladium or platinum.

Thus the reaction I'm thinking of would be:

2Ar-OH(l) + TiH2-x(s) => 2C6H6(g) + TiO2(s) + (1-x/2)H2(g) + heat

This reaction could be very aggressive (if it proceeds at all) and take note H2 produced, so be careful if you try it. For one thing, no ignition source anywhere near the outlet :o

I'm hoping this could similar to the chlorothermal reactions, start it with a little heat (but beware; H2 formed!) after which it should be self sustaining, almost aggressive.

I think a setup with condenser here is a must, for one thing to lead the hydrogen away from the source producing it. At least no chance of plugging it with solids this time ;)

Also, please be aware, hot phenol gives TERRIBLY nasty burns if it gets on you. REALLY take this threat seriously! If your apparatus cracks or pops and splatters the phenol and it gets on you face... well I've seen what that does. Needless to say, a $10 plastic face shield (not just goggles), gloves and long sleeve plastic overcoat makes ALL the difference when working with phenols. Please implement these, a technician in a lab I was working in was badly injured with a phenol experiment before implementing this basic safety protocol.

Dratz, I've gone and rambled too much again, kinda useless as you probably don't have titanium hydride powder :(

[Edited on 6-5-2014 by deltaH]

blogfast25 - 6-5-2014 at 04:47

DeltaH:

Not sure about TiH2 and as you say: I don't have any. It would become rather expensive way to prepare small amounts of benzene though...


deltaH - 6-5-2014 at 05:06

Yes agreed and as I feared :(

[Edited on 6-5-2014 by deltaH]

blogfast25 - 7-5-2014 at 04:07

Other methods for benzene preparation have been proposed here:

http://www.sciencemadness.org/talk/viewthread.php?tid=4902

No one seems to have followed up on anything other than benzoate/hydroxide though.

blogfast25 - 11-5-2014 at 09:49

Just finished a quick test tube test on Al + phenol.

A quarter of a teaspoon of phenol was loaded into a test tube and heated to BP. A long, thin strip of aluminium was then lowered into the boiling phenol. No reaction was observed. No discernible damage was sustained by the tip of the aluminium.

plante1999 - 11-5-2014 at 10:35

amalgamated aluminium should prove to be reactive in such a medium.

deltaH - 13-5-2014 at 07:55

I was reading the Wikipedia article on 'decarboxylation' and found this oddity at the bottom:

Quote:
In beverages stored for long periods, very small amounts of benzene may form from benzoic acid by decarboxylation catalyzed by the presence of vitamin C[8]


[8] http://www.cfsan.fda.gov/~dms/benzdata.html

Unfortunately, this is a dead link. While not practical, I was merely curious how vitamin C catalyses the decarbonylation of benzoic acid to benzene at near room temperature -- even in "very small amounts"!?

plante1999 - 13-5-2014 at 08:19

@blogfast

Amalgamated aluminium will also work when not into aqueus solution is used as an oxygen reductor, and not an hydrogen one needed to reduce alkanes or nitro group.

Nicodem - 15-5-2014 at 10:31

Quote: Originally posted by blogfast25  
I really don't know and only go by a reference in wiki.

The scientific literature is there to be used when doing scientific experiments. Reviewing the literature is the most important part of designing the experiment.

Prof. Baeyer: Ueber die Reduction aromatischer Verbindungen mittelst Zinkstaub
Journal für Praktische Chemie, 1867, 100(1), 46–47.
DOI: 10.1002/prac.18671000110

S. Marasse: Untersuchungen über das rheinische Buchenholztheerkreosot
Justus Liebigs Annalen der Chemie, 1869, 152(1), 59–87.
DOI: 10.1002/jlac.18691520105

E. Müller: Ueber pyrogene Zersetzung von Gasöl, Phenol und Kreosot
Journal für Praktische Chemie, 1898, 58(1), pages 1–50.
DOI: 10.1002/prac.18980580101

See also the review in DOI: 10.1002/ange.19290421404.

Attachment: 10.1002jlac.18691520105.pdf (1.1MB)
This file has been downloaded 440 times

Attachment: 18661400306_ftp.pdf (79kB)
This file has been downloaded 472 times

organicchemist25 - 18-5-2014 at 14:30

I have only made benzene by dry distillation of Benzoic acid and Sodium Hydroxide. It is simple to me. I bought a threaded pipe on both ends, bought two end caps and tapped one cap to fit a 1/4" compression fitting with about a 3' 1/4" copper line. Load up the vessel and flame that thing. I also took the burner part off of a portable deep fryer and connected it to a propane tank. I get approximately 250 mL of benzene. set up and distill and get very reasonably pure benzene. Simple.

[Edited on 18-5-2014 by organicchemist25]

aga - 20-5-2014 at 13:55

Would it be insanely stupid to try an aluminium vessel/setup rather than glass ?

[Edited on 20-5-2014 by aga]

gdflp - 20-5-2014 at 14:14

Sodium Hydroxide + Aluminum = Yes that would be stupid to try. You will most likely end up with benzene vapor/liquid going everywhere.

Organikum - 18-6-2014 at 10:54

Oh that was a long long time ago....lol

Found an article which fits in snuggly here I thought an here it is, hope it wasnt already posted, a quick search showed nothing though.

The guys actually wanted to make benzoic acid from CO2 and benzene but to get the direction they researched the decomposition of benzoic acid first and catalysts which can be used.
In short: Zn-Cu-Cr oxide catalyst gives 95% benzene in 100 hours at 290°C to 300°C, with additional MnO2 85% in 10 hours.

What might PAA and GAA do under the conditions. Hmmm...

/ORG



Attachment: benzene-from-benzoic-kinney1931.pdf (250kB)
This file has been downloaded 419 times


p2e3r4f5e6c7t8 - 1-10-2014 at 06:14

has anyone tried this method.
http://m.youtube.com/?#/watch?v=S8rtyRnZZMU

Sodium benzoate + Sodium hydroxide heated to decompose and vapour condensed & collected ,then distilled & dried.

Bert - 1-10-2014 at 08:11

Quote: Originally posted by p2e3r4f5e6c7t8  
has anyone tried this method.
http://m.youtube.com/?#/watch?v=S8rtyRnZZMU

Sodium benzoate + Sodium hydroxide heated to decompose and vapour condensed & collected ,then distilled & dried.


Yes. This method is featured on several other threads. The improvised equipment and dry distillation method shown in that video is functional.

p2e3r4f5e6c7t8 - 2-10-2014 at 01:58

So why then are newbees fuckin a round with soke of these more complex methods when this method is easy and dosent require any fancy glassware ?

some simple distilation glassware is required, and that just makes things a little easyer.
when i first started out intersted in chemistry it was for all the wrong reasons untill i learned the hard way (goal), Now its the case of all those wierd funny looking symbols and the like that looks intersting ,& the fact that you can make something from nothing so to speak.
Anyway back on topic , benzene is soooo simple to make and only requires otc chems to synthesis it from scratch witch is the best way to learn about the reaction.

Bert - 2-10-2014 at 07:11

Quote: Originally posted by p2e3r4f5e6c7t8  
So why then are newbees fuckin a round with soke of these more complex methods when this method is easy and dosent require any fancy glassware ?


Organic chemistry presents a LOT of possible paths to a target substance...

If the goal is PRODUCTION, for some other use- The quickest/simplest/cheapest path is the obvious choice.

If you are interested in learning as much as possible and acquiring skills in a wide variety of techniques, you might take a more indirect path. Just for the practice in lab technique, and to validate your grasp of theory.

Most people here are interested in gaining knowledge and proficiency?

CuReUS - 3-10-2014 at 23:48

Quote: Originally posted by Bert  




If you are interested in learning as much as possible and acquiring skills in a wide variety of techniques, you might take a more indirect path. Just for the practice in lab technique, and to validate your grasp of theory.

Most people here are interested in gaining knowledge and proficiency?


absolutely correct
by trying indirect routes ,chemists learn about new reagents,new reactions ,and new rearrangements;) (for example,see synthesis of vitamin B-12)

by the way, why cant we try fenton's reagent(H2O2 + FeSO4) on benzoic acid to make it benzene

p2e3r4f5e6c7t8 - 6-10-2014 at 04:08

hmm, you make a good point bert.
And that is also very true cureus

CuReUS - 6-10-2014 at 04:26

Quote: Originally posted by p2e3r4f5e6c7t8  

And that is also very true cureus


but i dont think we could use fenton's reagent directly,it is too powerful an oxidising agent and might break the ring(it is used to clean glassware and remove stains) .we might have to alter the proportion of peroxide to Fe ions for doing the decarboxylation;)

S.C. Wack - 1-11-2014 at 09:54

Quote: Originally posted by HgDinis25  
http://web.anl.gov/PCS/acsfuel/preprint%20archive/Files/48_2...


Britt and Buchanan published a later version of the above, attached, and other similar things such as
DECARBOXYLATION OF SALTS OF AROMATIC CARBOXYLIC ACIDS AND THEIR ROLE IN CROSS-LINKING REACTIONS
http://web.ornl.gov/~webworks/cppr/y2000/pres/109540.pdf

These all lead to perhaps endless somewhat-interesting-DOE-reference chasing, e.g. Manion who does somewhat interesting things in naphthalene with substituted benzoic acids etc., and/or decarboxylation agents in this same journal; probably no one ever uses NaOH and benzoate in a soup can though.

Note that the attachment's "Commercial Calcium Benzoate" (7% water) at 500C produced almost pure benzene, while "Vacuum-Dried Commercial Calcium Benzoate" and dried "Synthesized Calcium Benzoate" gave mostly benzophenone, but in much lower yield, in line with the 30% yield of benzophenone from the dried calcium salt in Systematic Organic Chemistry. The quoted article mentions that water is very favorable for producing benzene from sodium benzoate, for whatever reason(s).

Attachment: ef.pdf (244kB)
This file has been downloaded 341 times

BromicAcid - 1-11-2014 at 11:08

Quote: Originally posted by CuReUS  
by the way, why cant we try fenton's reagent(H2O2 + FeSO4) on benzoic acid to make it benzene


Fenton's reagent can add -OH groups directly to the ring via radical reaction, you'll end up with a mix of phenols and such.

CuReUS - 2-11-2014 at 05:59

Quote: Originally posted by BromicAcid  
Quote: Originally posted by CuReUS  
by the way, why cant we try fenton's reagent(H2O2 + FeSO4) on benzoic acid to make it benzene


Fenton's reagent can add -OH groups directly to the ring via radical reaction, you'll end up with a mix of phenols and such.


yes you are right ,but the yield of phenols is very small .i only suggested the idea because fenton's reagent is used in oxidation of sugar and suppose you have an aldose ,it oxidises the terminal aldehyde to carbon dioxide and converts the next secondary alcohol in the chain(6 carbons in a chain) to aldehyde

btw,as fenton's reagent produces hydroxyl free radicles and the mechanism of NBS allylic bromination is a free radicle one,can we do allylic hydroxylation using fenton's reagent

palladium8 - 29-11-2014 at 20:41

Could the benzoate to benzene be done in a microwave?

I'm getting started with microwave chem and wondered why this couldn't be applied. Sodium hydroxide and sodium benzoate will both absorb MW. It will solve the problem organikum got, the sticky layer which was caused by uneven heating.

Perhaps the MW can do a whole lot of other high temp destructive distillations. I really don't want to mess with flames and that's the only other reasonable way.

subsecret - 27-12-2014 at 10:24

I noticed that someone posted a reference where benzene was produced from toluene, HCl, and AlCl3, which causes some of the methyl groups to migrate between benzene rings, forming xylene, mesitylene, etc.

Wouldn't a longer reflux time work to produce benzene? After all, if this is the reverse of a Friedel-Crafts, chloromethane should be produced, which would boil off at a low temperature. This should drive the equilibrium to the side with the benzene.

I didn't see the original PDF, so here it again for consideration.


Attachment: The Friedel-Crafts reaction Part III.pdf (287kB)
This file has been downloaded 682 times


aga - 1-1-2015 at 11:11

After a few attempts at synthesising Benzene from PVC pyrolysis, i'm giving up.

Not one experiment yielded any useable benzene.

You mostly get HCl, Carbon, and a small amount of brown gunk that makes a huge mess of the glassware.

CuReUS - 1-1-2015 at 21:10

Quote: Originally posted by aga  

You mostly get HCl, Carbon, and a small amount of brown gunk that makes a huge mess of the glassware.


what about the large number of carcinogenic chemicals like dioxins etc. :o

but no one tried the method posted by this guy,maybe it was too crazy
http://www.sciencemadness.org/talk/viewthread.php?tid=508

[Edited on 2-1-2015 by CuReUS]

subsecret - 1-1-2015 at 21:49

Quote: Originally posted by CuReUS  

...
what about the large number of carcinogenic chemicals like dioxans etc. :o
...


As stated by aga, the PVC method isn't practical. But while on the topic of dioxins, thermolysis may be a good way to get rid of these compounds. Run the gas stream into the bottom of a hot fire to decompose any undesirable waste.

UC235 - 2-1-2015 at 02:06

Quote: Originally posted by Awesomeness  
Quote: Originally posted by CuReUS  

...
what about the large number of carcinogenic chemicals like dioxans etc. :o
...


As stated by aga, the PVC method isn't practical. But while on the topic of dioxins, thermolysis may be a good way to get rid of these compounds. Run the gas stream into the bottom of a hot fire to decompose any undesirable waste.


Except that "backyard" burning of organochlorine containing waste is the single largest source of environmental dioxins. They aren't necessarily formed as the material burns and more fire won't fix it. With regards to incinerators (formerly the major source), wiki has this to say:

"In incineration, dioxins can also reform or form de novo in the atmosphere above the stack as the exhaust gases cool through a temperature window of 600 to 200 °C. The most common method of reducing the quantity of dioxins reforming or forming de novo is through rapid (30 millisecond) quenching of the exhaust gases through that 400 °C window. Incinerator emissions of dioxins have been reduced by over 90%"

maleic - 13-1-2015 at 01:00

Propylene and benzene under the catalysis of aluminium trichloride, from 80 to 90 degrees Celsius for hydrocarbon reaction, get the isopropyl benzene. Isopropyl benzene with air in the 100-120 degrees Celsius, and 100-400 kpa pressure isopropyl benzene oxidation generating hydrogen peroxide. Isopropyl benzene with sulfuric acid hydrogen peroxide under 60 ℃ atmospheric pyrolysis for phenol and acetone.

subsecret - 10-5-2015 at 17:26

Just a note:

JB Weld will does not withstand the effects of dry distillation of sodium benzoate and sodium hydroxide. I sealed a copper tube into a paint can with it, but several minutes after I applied heat, it began to flake and leak benzene vapor.

DFliyerz - 29-7-2015 at 07:54

I made benzene with sodium benzoate and sodium hydroxide too, and I think I have an idea of what causes the orange/yellow colors. After researching for a little while, it seems that the yellow/gold could be caused by chrysene, and the orange caused by tetracene.

maleic - 8-9-2015 at 23:56

Sodium benzoate in soda lime heating conditions can generate sodium carbonate and calcium carbonate and benzene, due to the heating condition, so the temperature is not very high, so the sodium carbonate and calcium carbonate does not decompose, and the boiling point of benzene is 353.25 K (80.1 DEG C), so the benzene will be volatile, so chemical balance will be broken, and toward the direction of benzene formation reaction. Benzene vapour is cooled and condensed, so that you can collect benzene.

My Recent Attempts:

ScienceHideout - 15-10-2015 at 18:13

Hi guys!

This week I tried three different possible routes to benzene. They all failed, but are worth noting.

Firstly, I know some of you are bound to think that I have been sniffing too many of my reagents. I promise you that I am sane. The most likely reason why you would be led to such a hypothesis is because of the fact that I am using rather odd reagents to do such reactions. The reason is that I prefer not working with mixtures of ungodly hot hydroxides and benzoic acid... just something I prefer to stay away from. So here, as follows, are my recent experiments, and also one of my ideas ;)

1) REDUCTIVE DE-AMINATION A couple days ago I wanted to try to make it from aniline. I took about 2 mL of aniline and added to it a stoiciometrically correct amount of HCl to produce a required amount of HNO2 when a nitrite is added in situ. I then made the nitrite solution and added it to get a diazonium salt. After the salt was made, I tried reducing it with a HUGE excess of NaBH4. Borohydrides have four working hydrides, so I figured a fourth of an equivalent should work, but I doubled it and chose to do a 2:1 molar ratio of diazonium salt to NaBH4. As I added it, it made a fizzing sound but no other change was noted. I got, instead, a black tar in the reaction vial and it didn't smell like benzene, so I disposed of it. The tar was soluble in most polar organic solvents, though.

What went wrong?: I considered for a moment that the BH4- might have been reacting with the water instead of the diazonium, but I know that borohydrides react with water slower than it reacts with other chemicals. Maybe a hydride just won't react well with this, I can't find a mechanism. Plus, see my note for the next step.

2) In this experiment, conducted the first steps very similarly (on a somewhat larger scale, .1 mol) but I tried warming the diazonium salt with EtOH instead of borohydride. I noticed the tar forming again, but in addition I saw many more bubbles than the previous reaction, so I thought it was working. Cool oily patterns were forming on the surface of the liquid. I dumped it in a sep funnel after it was done bubbling, but saw no separation like I expected. Instead, the crappy oil came out dropwise and made clean-up a disaster.

What went wrong?: I have no idea. It is either me or my reagents. The bottle of aniline I have is VERY old and has a sort of brown tinge... It also doesn't smell as pungently as I imagined it would (I inherited this chemical, it came from Merck originally. Must be about 60+ years old. Is there any way to test it and see if it is still good?

3) ZINC REDUCTION. Here I took .05 mol of phenol and melted it in a 5 mL reaction vial. I added .05 mol of zinc and attached a hickmann still head with a condenser at the top. I ramped up the heat until it started boiling and collected the distillate. I stopped the distillation just before it ran dry. I then took it apart and got the distillate. It crystallized. I just distilled the phenol. Not what I was after, but if anyone wants the purest phenol in town you are more than welcome to come to my house and scrape it out of the head!

What went wrong?: Maybe I should've refluxed it? I really don't know. A couple sources say it should've worked

WHAT'S NEXT? How about BENZENE from DOLLAR STORE CHEMICALS!
All you need is some para-dichlorobenzene mothballs, magnesium, ether from engine starting fluid and water! Tell me if you guys think it will work, I attached the mechanism, but I am very terrible at grignard-type radical mechanisms so please let give me your critiques!



mechanism.jpg - 110kB

Morkva - 16-10-2015 at 09:00

1. Aniline.

If there was an excess of aniline, could it be diazo tars?

Also, copper might help. The N2 needs to come off somehow: copper salt gives an electron and the nitrogen leaves the aryl radical. If you have copper salt and borohydride in solution I do believe there is at least transient Cu (I) hydride, which may be your key to success, if the issue was not just separation. Traditionally hypophosphite was used for this kind of reduction, also by a radical mechanism. This kind of radical chemistry is very interesting to search for and learn about.

2. Phenol. Zinc dust? Maybe if you ran the phenol vapors through a heated tube filled with zinc dust. Nobody can blame for not feeling like doing such a thing.

3. PDCB
I shall probably eat my words but you may not have to separately form the Grignard reagent or even use ether. Depending on the fondness you might get away with magnesium metal and alcohol. At least, I saw something like that with calcium metal "arene dechlorination", but then why would they use calcium? It is also probably possible to take off only one chlorine.

http://archneur.jamanetwork.com/article.aspx?articleid=17887...

In the Octopuses garden...



[Edited on 16-10-2015 by Morkva]

ScienceHideout - 16-10-2015 at 09:28

Thanks for the response, Morkva!

1) The copper sounds like it would be a good idea, I am just confused as to how I can get those ions in there without introducing any Br- or Cl- that would cause just a normal sandmeyer reaction. I find it interesting that you bring up the hypophosphite, because every source online uses phosphorous acid to do the reaction, which is very hard to obtain. I was unaware that it involves a radical mechanism, though. I thought it would be polar. Perhaps this is why I couldn't figure out how to draw the darn thing :P

2) I found many sources that describe passing vapors over red hot zinc dust, but a couple more questionable ones about plain distillation over zinc. Chemists are lazy; I just had to try the easier (and more practical) of the two! I assume that this is some sort of radical reaction, as well. Maybe the transition state is too high in energy? In which case, maybe an added catalyst would work.

3) Thanks for the [somewhat comical] link! I will study this! What you said, though, about not having to form grignard reagents in ether or anything though seems so... simple! I would've thought if it was that easy, others would've already done it...

Morkva - 17-10-2015 at 07:31

https://archive.org/stream/aromaticdiazocom031270mbp/aromati...

Look at p. 271 and on, it's about equidistant from the top.
On p. 282 it claims copper hydride works to halogenated by forming cuprous halides, so you're right, you would have to separate the ions and I do not see how either.

http://www.researchgate.net/profile/Cristian_Simion2/publica...

Phenyl chloride is claimed to be formed quantitatively, and was not reduced further, from dichlorobenzene. Ethanol was the most effective alcohol they tried. I have no idea about magnesium, except that is too an alkaline earth metal...

If CuH is still interesting to anyone.

http://pubs.acs.org/doi/pdf/10.1021/ic5027009

Open access article on the surface structure of copper hydride produced by different methods. One of these is by Cu(ii) and BH4 in aqua, which forms a coffee colored precipitate different in color from Würtz's old fashioned CuH which made with hypo phosphorous acid and is red.

Apparently copper hydride can decompose explosively when dried. It also decomposes slowly when wet.

http://journals.iucr.org/b/services/forthcoming.html
Structure and spectroscopy of CuH prepared via borohydride reduction
Elliot L. Bennett, Thomas Wilson, Patrick J. Murphy, Keith Refson, Alex C. Hannon, Silvia Imberti, Samantha K. Callear, Gregory A. Chass and Stewart F. Parker

We show by a combination of diffraction and spectroscopic methods that CuH produced by borohydride reduction of a CuII salt consists of a core of CuH with the Wurtzite structure and a shell of water. We also demonstrate that the shell is exchangeable for ethanol.


[Edited on 17-10-2015 by Morkva]

S.C. Wack - 17-10-2015 at 10:00

It would not have been unwise to purify the aniline before doing anything with it; of course you know that at the diazo stage, sodium stannite (NaOH and SnCl2) could have substituted for NaBH4, or the direction could have branched off towards phenylhydrazine with sulfite...which gives benzene with CuSO4.

clearly_not_atara - 14-11-2015 at 21:29

http://en.wikipedia.org/wiki/Benzene_in_soft_drinks

If this reaction occurs in cold soda cans it's probably much faster at high temperatures with higher catalyst loading (more Fe2+/Cu2+ and more ascorbate). I'm pretty sure that if you can't get benzoate for some reason your next best bet is to oxidize toluene or benzyl alcohol.

Dichlorobenzene will be quite slow to eliminate with magnesium, and probably won't react with zinc or lithium. It could be done by adding dichlorobenzene to a solution containing ethylmagnesium bromide. In any case it sounds hard :/

Terephthalic acid should decarboxylate as easily as benzoic acid and can be obtained from PETE.

[Edited on 15-11-2015 by clearly_not_atara]

Heavy Walter - 17-12-2015 at 05:51

Hi
Just joined the forum, I learn some people was interested in benzene synthesis.
I experimented with catalyst and acetylene, trimerizing it in order to prepare deuterated benzene. Reaction was quantitative and everything proceeded as expected, except the small acetylene generator I built, because gas generation was too violent. So I decided to freeze the D2O, landing a small chunk of calcium carbide on it, closing the generator and then I was able to control the pressure on it opening the way to a storage balloon.
The basic circuit is as follows: gas generator (like a calcimeter), three cold traps (ice, ice & acetone, liquid nitrogen), a storage balloon, one reactor with the catalyst and a cold finger were the C6D6 was collected while the catalyst was heated.
All the line was evacuated with rough pump and diffusion pump to 10-6 torr.
Regards!

brubei - 13-3-2016 at 16:43

Thx Walter ! That sound very efficient.

For those who tried with diazonium, hypophosphorus acid was the right reducing agent. It seems to be the classical method for reducing aromatic amines

http://chemwiki.ucdavis.edu/Core/Organic_Chemistry/Amines/Re...

One would probably try with other H• radical donor

[Edited on 14-3-2016 by brubei]

[Edited on 14-3-2016 by brubei]

vin123 - 13-3-2016 at 23:14

Depending on the purpose an amount of your preparation of benzene (solvent or building block), in the latter case you may try arenediazonium salt reduction (starting from aniline). It is possible to trap them with tetrafluoroborate salts to become more stable but I would use them in situ. But be very carefull when scaling up. For solvent use I would try something else. Moreover, their are many replacements for benzene as a solvent, why using it if its not available?

aga - 26-11-2016 at 10:53

Rain today, so tried out the benzene synthesis by dry distillation of 50g sodium benzoate & 14g sodium hydroxide.

(Benzene is the last item required to attempt Magpie's Congo Red synthesis)

Initially a 'golden syrup' tin was used as the reaction vessel, with the remaining distillation glassware (the rest was in the wash).

rig1.JPG - 61kB

Using just a small spirit burner, this actually worked (kind of) and gave a tiny, yet totally clear product which smelt distinctly like i always thought benzene smelled.

1styield.JPG - 55kB

A blowtorch was then used to raise the temperature and speed up the reaction. Perhaps doing that caused the next distillate to become orange ?

Unfortunately the metal tin seal began to leak, also the lid, which led to to small benzene fire which was immediately extinguished.

Undeterred, the contents of the tin were extracted and around half was loaded into a 100ml RBF. They're the cheaper to replace than the 250ml+ sizes.

rig2.JPG - 54kB

Again the spirit burner was used, this time with aluminium foil for insulation.

The product came over after 8 minutes at a rate of 1 drop every 3 seconds.

The gas escaping from the pot appeared as 'threads' in the condenser.

gasthreads.JPG - 40kB

The result of this distillation was a yellow/orange liquid which clearly has two layers.

2layer.JPG - 46kB

The rest of the reagents will be distilled soon, the products combined, then re-distilled to work out the yield.

[Edited on 26-11-2016 by aga]

aga - 27-11-2016 at 14:29

This all went a bit wrong and destroyed a 100ml and a 250ml RBF on cooling, as expected.

All of the available sodium benzoate was used (~80g) resulting in about 37ml of orange liquid, still to be re-distilled to get any actual product.

The spread of heat appears to be vitally important.

In an RBF one can see the area affected by the flame becoming black and bubbling, while the rest of the white powder remains unaffected.

With a naked flame (spirit burner) this can be moved to target an unaffected area, which seemed to squeeze out a brief run of more drops of distillate each time the flame was moved.

A steel conical flask with a huge bottom surface area might work, so that the powder remains in a very thin layer and gets heated enough for the reaction to complete.

Maybe a steel reactor where an inert atmosphere is super-hot and the powder gets fed in somehow, while allowing the benzene gas to escape.

Edit:

The products were redistilled and a cloudy liquid came over between 71 C and 75 C, the orange stuff stayed in the boiling pot.

Yield at this point is 24.1g = 56.7%

Probably needs drying so that % will reduce :(

[Edited on 28-11-2016 by aga]

sulfuric acid is the king - 26-2-2017 at 15:55

What about coal pyrolisis?...

macckone - 26-2-2017 at 20:41

Coal pyrolysis is the traditional method but it requires
huge amounts to coal.

JJay - 26-2-2017 at 22:22

Quote: Originally posted by aga  
This all went a bit wrong and destroyed a 100ml and a 250ml RBF on cooling, as expected.

All of the available sodium benzoate was used (~80g) resulting in about 37ml of orange liquid, still to be re-distilled to get any actual product.

The spread of heat appears to be vitally important.

In an RBF one can see the area affected by the flame becoming black and bubbling, while the rest of the white powder remains unaffected.

With a naked flame (spirit burner) this can be moved to target an unaffected area, which seemed to squeeze out a brief run of more drops of distillate each time the flame was moved.

A steel conical flask with a huge bottom surface area might work, so that the powder remains in a very thin layer and gets heated enough for the reaction to complete.

Maybe a steel reactor where an inert atmosphere is super-hot and the powder gets fed in somehow, while allowing the benzene gas to escape.

Edit:

The products were redistilled and a cloudy liquid came over between 71 C and 75 C, the orange stuff stayed in the boiling pot.

Yield at this point is 24.1g = 56.7%

Probably needs drying so that % will reduce :(

[Edited on 28-11-2016 by aga]


I made my benzene in steel solvent cans. I punched a hole in the top and attached a brass compression fitting, which I connected to a copper tube that ran through a thermometer adapter into a condenser. I loaded a powdered mixture of sodium hydroxide and sodium benzoate into the cans, and heated I heated the cans over a propane stove. This resulted in an orange distillate, which was washed, dried, and fractionated. Each can is good for several runs.

The yield was far from quantitative, but the reaction is easy to do, and the reactants are cheap and readily available. I was very careful not to breathe any benzene; I'm not really sure how severe a carcinogen it is, but it has been proven to actually cause certain cancers.



[Edited on 27-2-2017 by JJay]

The jersey rebel - 30-4-2017 at 16:34

The methods I am aware of that are accessible to the amateur are demethylation of methylbenzenes, oxidation of toluene to benzoic acid and decarboxylate the acid over copper chromite or sodium hydroxide, if acetylene is readily available it can be catalytically cyclotrimerized to form benzene with a cyclotetramer side reaction, Decarboxylation of terephthalic acid over a catalyst such as basic copper carbonate or copper chromite, acetone trimerization to form the methylated threefold ketone of benzene, and deoxygenation of phenol with powdered zinc. There are numerous other methods though the feasibility is entirely contingent on where you live and your equipment/experience. These methods are for those who happened to live on the east cost of the US. Accessibility will differ from country to country and even on the state level.

plastics - 5-5-2017 at 03:58

Quote: Originally posted by JJay  
Quote: Originally posted by aga  
This all went a bit wrong and destroyed a 100ml and a 250ml RBF on cooling, as expected.

All of the available sodium benzoate was used (~80g) resulting in about 37ml of orange liquid, still to be re-distilled to get any actual product.

The spread of heat appears to be vitally important.

In an RBF one can see the area affected by the flame becoming black and bubbling, while the rest of the white powder remains unaffected.

With a naked flame (spirit burner) this can be moved to target an unaffected area, which seemed to squeeze out a brief run of more drops of distillate each time the flame was moved.

A steel conical flask with a huge bottom surface area might work, so that the powder remains in a very thin layer and gets heated enough for the reaction to complete.

Maybe a steel reactor where an inert atmosphere is super-hot and the powder gets fed in somehow, while allowing the benzene gas to escape.

Edit:

The products were redistilled and a cloudy liquid came over between 71 C and 75 C, the orange stuff stayed in the boiling pot.

Yield at this point is 24.1g = 56.7%

Probably needs drying so that % will reduce :(

[Edited on 28-11-2016 by aga]


I made my benzene in steel solvent cans. I punched a hole in the top and attached a brass compression fitting, which I connected to a copper tube that ran through a thermometer adapter into a condenser. I loaded a powdered mixture of sodium hydroxide and sodium benzoate into the cans, and heated I heated the cans over a propane stove. This resulted in an orange distillate, which was washed, dried, and fractionated. Each can is good for several runs.

The yield was far from quantitative, but the reaction is easy to do, and the reactants are cheap and readily available. I was very careful not to breathe any benzene; I'm not really sure how severe a carcinogen it is, but it has been proven to actually cause certain cancers.

[Edited on 27-2-2017 by JJay]


Exactly the way I did it. Stainless biscuit barrel, tank connector and some 15/22mm copper pipe plus fittings. Yield not great but better than nothing. I got the same orange distillate which I then redistilled in a glass setup


B9734456-AC04-445C-B82D-12B878481D79.JPG - 1.1MB

 Pages:  1    3