Sciencemadness Discussion Board - Preparation and uses of Organic Carbonates

Preparation and uses of Organic Carbonates

BromicAcid - 3-12-2004 at 10:52

After a Eclectic mentioned an interesting patent in the thread on the preparation of phosgene from carbon tetrachloride I decided to give it a try. It is currently simmering following this procedure (Patent EP0638541:

Quote:

A 300-ml three necked flask equipped with a Snyder's fractioning column having a reflux condenser in an upper portion and a thermometer was charged with 60.1 g (1.00 mol) of urea, 77.6 g (1.25 mol) of ethylene glycol and 3 g of zinc oxide, and while the mixture stirred, the pressure in the flask was reduced to 100 mmHg and the mixture was heated to 145C. The mixture was allowed to react for 2 hour, and then the reaction mixture was cooled. The amount of the reaction mixture was 106.1 g. The reaction mixture was analyzed for a composition by gas chromatography to show 15.4 g of unreacted ethylene glycol and 86.0 g of formed ethylene carbonate.

The above results show the following. The conversion of ethylene glycol is 80.1 % under the theoretical value of 80%, the selectivity to ethylene carbonate on the basis of the reacted ethylene glycol is 97.5% and the selectivity to ethylene carbonate on the basis of urea is 97.6% (conversion of urea = 100%).

But of course there were exceptions, I used a 500 ml one neck round bottom. The thermometer was in the oil bath around the flask and I skipped the refluxing column. As such the only neck was fitted with a gas exit tube running to a air conditioner compressor. It doesn't quite pull what it needs to but I'm just going to let it run with what it's got. The reaction is running smoothly and even when the flask is removed from the oil bath it constantly bubbles from ammonia evolution. The exit gasses are being run though dilute H2SO4 and the compressor appears to be leaking tetramine copper complex ... more updates will come when I am done. I am hopping the ethylene carbonate crystallizes out but if it doesn't I may well convert it to methyl or ethyl carbonate which is an easy conversion.

The organic carbonates are becoming widely used solvents due to their low toxicity. And somewhat low reactivity. Their widest application as of now is in the preparations of poly carbonates. Solvents in the plastic industry in general and solvents in lithium batteries.

Update: There are wispy white fumes that appear right where the gasses are being removed from the flask. The mixture has 45 more minutes to heat. The volume of liquid appears to have decreased slightly. When I first started this there appeared to be a greater volume of urea then ethylene glycol but it all become homogenous shortly. I used antifreeze for my ethylene glycol. Aside from the green colorant the ingredients were listed as ethylene glycol (95 - 98%) and diethylene glycol (2 - 5%) so nothing that should disrupt the process to any notable extent. I did forget to add the correct amount of zinc oxide though, I only added 2 grams instead of 3 hopefully that will not hurt the reaction too badly.

Update #2: I ended up with a solution that looked almost identical to orange juice except it was thicker. My initial weight of reagents was ~142g and my final weight of solution was ~123 g therefore according to:

HOCH2CH2OH + NH2CONH2 ---> (-OCH2CH2OCO-) + 2NH3

Assuming my exit gasses were entirely ammonia and they constituted only products from the reaction above, only a little over half the urea present reacted. The resulting solution smelled strongly of ammonia. Nothing precipitated out of it except the fine powder of ZnO which I did not include in the above. Reasons for the reduced yield might include 1) Insufficient vacuum 2) Less then desired amount of catalyst 3) Contaminates in antifreeze may have been detrimental

However my bet would be with the first and second of those. I will do more analysis of my solution later. Things didn't really work out badly at all, through out the reaction the temperature stayed between 143 and 155 C. I may try this again when the weather gets warmer. Fractional distillation is always an option to attempt to separate the reactants, however both ethylene carbonate and ethylene glycol both have high boiling points and it makes it sound unattractive. I could convert to the methyl ester though and attempt to distill that (Bp 90.6 C).

[Edited on 12/3/2004 by BromicAcid]

chemoleo - 3-12-2004 at 13:17

Very interesting!

A couple of questions:
I take it ethylene carbonate only exists as the ring form, and not as polymerised ethylene glycol -carbonate chains?
Also, what is the bp of it that you mention it would be hard to distill it off? Is it soluble in water (I guess due to its polar nature this is to be expected)?
I take it the low pressure is chosen so as to rip the nascent NH3 out of the system, and to shift the equilibrium of the reaction to the right side? Presumably that in part is the reason why the weight loss is not as high as desired.

Aside from making low carbonates, such as ethyl/methyl, shouldnt the ethylenecarbonate be a useful precursor to heterocyclics? What else?

Anyway, isn't it awesome that a reaction that ordinarily'd require phosgen works with simple ingredients such as urea and glycol?

BromicAcid - 3-12-2004 at 15:16

I would have thought that this procedure would have yielded a polycarbonate at least as a competing side product, but the patent says that the product is almost entirely ethylene carbonate and I got no evidence of polymerization (e.g., an insoluble product or strange foaming). There is probably some sort of intermediate formed whereby the tendency is towards cyclisation rather then condensation with another ethylene glycol molecule. I have been thinking about this mechanism and it is quite an interesting puzzle.

As for the boiling points Ethylene carbonate has a boiling point of 248C and ethylene glycol of 197.2, should be possible to separate the ethylene glycol from the mixture the boiling points are dissimilar enough but with the other impurities in the antifreeze the ethylene carbonate left in the boiling flask may well contain a number of impurities. Reduced pressure distillation may be of benefit here. But the diethylene glycol impurity has a Bp of 245 so unless I treat it chemically it might be impossible to separate by distillation.

And for the solubility of ethylene carbonate look at this list

Quote:

Mixcible (40%) with water, alcohol, ethyl acetate, benzene, and chloroform; soluble in ether, n-butanol, and carbon tetrachloride

Wide solubility here....

I'm guessing the low pressure is to rip the NH3 out of the system just as you think too, possibly to decompose some sort intermediate. I would like to run this reaction for a while and weigh the reaction flask every hour to monitor the weight loss. So I can find out when the reaction is most complete. I wonder how a large excess of catalyst would affect the reaction.

I'm wondering too about what kind of synthetic reaction might be afforded with carbonates, under each carbonate entry under uses it states 'organic synthesis' but I've never really seen a reaction that calls for the use of ethylene carbonate. Ethylene carbonate says it's used for hydroxyethylation reactions. I wonder how it would react with anhydrous acids and such, ethylene carbonate and concentrated H2SO4 might yield organic sulfates and such. If I can get this method to work good I would like to devote a large portion of time to working with carbonates as reagents and as solvents, there just seems to be so little known about them but the carbonate functional seems full of possibilities.

Eclectic - 31-12-2004 at 20:07

It probably wouldn't hurt to distill your ethylene glycol from the antifreeze first to get rid of silicates and phosphate corrosion inhibitors. I think if you used an excess of urea and ran the reaction at reflux (with an air cooled condenser),it would go to completion with the formation of cyanuric acid in ethylene carbonate and no vacuum required.

[Edited on 1-1-2005 by Eclectic]

BromicAcid - 9-1-2005 at 12:05

My reaction product that I obtained at the start of this thread has just recently precipitated something. It sat relatively homogenous and liquid for weeks after it was made despite the temperatures outside being below -5 C, however when I checked on it today about 3/4 of the mixture had solidifed at the bottom, forming a mass that looked like brain matter. The remaining mixture on top of it is still liquid. Before the weather heats up I will have to remove the liquid above and preform some experiments on the resulting solid at the bottom. But if they did separate, the product from the reactants (rather then it just being some homogenous solid), I will be considerably happier since high-temperature distillation will have been avoided.

sparkgap - 16-1-2005 at 05:33

I'm quoting first:

QUOTE from BromicAcid

...the product is almost entirely ethylene carbonate and I got no evidence of polymerization...

Quantitative yield, huh? My hunch on this is that the molecule would prefer to react with itself (alright, intramolecularly), rather than chain up with its nearest neighbor. Entropy-wise, that's a good thing. That's why the guys who make polycarbonate (Bayer, G.E., and a few others) use phosgene and not urea. Not too reactive.

Well, apart from being a good polar aprotic solvent, the only reaction I can think off my head for carbonates like them is Claisen's ester condensation. Then again, sodium ethoxide might not always be handy... right?

The thing is, ethylene carbonate can exhibit a good portion of the reactions esters can undergo, so I'm guessing that pairing it along with concentrated sulfuric acid will only afford you glycol and carbon dioxide. Wasted effort.

Just my two cents worth...

BromicAcid - 16-1-2005 at 18:45

Well, it would be neat to get organic sulfates from carbonates but if it doesn't appear to work, oh well.

I was looking up some information on ethylene carbonate the other day and photocopied some information (which I failed to write down the source of, though I remember which book it is) that may add to the meat of this thread.

For preperations it mentions bubbling CO2 through highly basified methanol to give a small precentage of methyl carbonate. Likewise direct carbonation of alkoxides (in the presence of selenium or Cu(OMe)2 gives organic carbonates). Phosgene reacting with alcohols produces an intermediate chloroformate which can be isolated so this could also serve as a starting point although it is none to avalible either. Heating N,N'-carbonyldi-imidazole (Which looks very much like urea

) with the alcohol of your choice gives organic carbonates. Also mentioned is the reactoin between carbonate anions with alkyl halides in DMF giving the esters. And finally it is mentioned that treating esters of dichloroacetic acid with sodium hydride gives the carbonate in 30 - 70% yield.

BromicAcid - 4-4-2005 at 14:03

My solid that crystalized out of the ethylene carbonate mixture re-liquified yesterday as the temperature outside neared 25C. I had poured off the liquid that failed to freeze while it was cold so I took the liquid I ended up with and put it in my freezer to re-freeze it. It froze readily and the soid I got was kind of like ice cream in consistency.

Obviously if it melted at 20C and it has a normal melting point of near 40C then there are some impurities present, how to purify from here?

If I can find a good method to get from here to a decent product, getting to here is just a matter of freezing the mixture in a freezer and pouring off the portion that fails to crystalize. Any ideas?

Picture attached.

frozen.jpg - 47kB

JohnWW - 4-4-2005 at 18:33

Phosgene, COCl2, is deadly poisonous, like CO, so if you are making carbonate esters at home by hydrolysis reactions with alcohols/phenols, whether monomeric or polymers, you had better be very careful. Wear a military-style gas mask.

Another point is:- why is sheet polycarbonate, usually 3 or 4.5 mm thick, available clear or tinted, as used for unbreakable window glazing, SO MUCH MORE EXPENSIVE than alternatives such as perspex (polymethylmethacrylate, which is not as strong) or ordinary plate glass? The materials used to make it, COCl2 and end-diols or diphenols, and the process, cannot be significantly any more expensive than those use to make perspex, which is made by the free-radical polymerization of methyl methacrylate. I am sure the patents on polycarbonates should have expired by now, as they have been around much longer than 20 years, so competition should have brought their prices right down. Or perhaps glazing-grade polycarbonates are made using secret materials and processes, which have never been patented, and are still "trade secrets"?

sparkgap - 5-4-2005 at 00:16

I don't think he's using phosgene though; looking at the previous posts, he's doing it using urea and glycol.

I think it's not the low cost of the raw materials per se that make polycarbonate expensive; rather, it may be that the raw materials, phosgene for instance, may require special handling. Then there is also post processing of the raw polymer for consumer use.

sparky (°..°)

BromicAcid - 5-4-2005 at 14:21

Although I origonally thought to make the ethylene carbonate using carbonyl chloride I decided against it when I found this prodcedure. Today I tried to extract my ethylene carbonate from the mixture, first with benzene, then acetonitrile, and finally chloroform, the chloroform made an emulsion or was simply miscible, it's sitting right now and has been for over and hour to see if it seperates. I will evaporate the benzene and acetonitrile solutions and see if the ethylene carbonate went into them to any extent.

Mybe I'll try eluding it through alumina or something. The original procedure reacted the mixture to make the ethylene carbonate and that was used directly to make methyl carbonate by reacting the resulting mixture with methanol under a few times atmospheric pressure so it can be heated to 130C and even then yields of the methyl ester are <40%. What am I to do?

There has been lots of work lately on using methyl carbonate and ethylene carbonate for many things being that they are green chemicals. They make good high temperature alkylating agents among other things and are just overall interesting.

sparkgap - 6-4-2005 at 23:54

Bromic, was the ethylene carbonate-chloroform mixture clear or on the cloudy side? If it was clear, even after long standing, I don't forsee any separation happening.

Same question goes for the other two solvents you used.

How fine is your alumina?

sparky (^_^)

BromicAcid - 7-4-2005 at 19:43

My chloroform solution did eventually separate, though the area between the two liquids is a little hazy. A teacher at my school suggested I just distill what I have, dispite ethylene glycols similar boiling point, if it separated like it did then there is a certain degree of purity to it already and hopefully the majority of the freezing point depressant is urea with only a small amount of ethylene glycol.

As for my alumina, I don't have any

I should buy some good activated alumina but I could use some calcium carbonate possibly, that is what they used in the early days when chromatography was first developed. I just need to get a good sample of ethylene carbonate so I can actually see how it behaves with certain solvents and such so I might have a easier job of separating it in the future.

BTW, I plan on running the reaction to produce the ethylene carbonate in a large pressure cooker with a large batch if I have any luck with a simple distillation.

BromicAcid - 9-5-2005 at 17:29

Found today that ethylene glycol forms an azeotrope with xylene that boils around 135 C and has a composition 87% Xylene 13% Ethylene glycol, so if Xylene does not from an azeotrope with ethylene carbonate then...

1) React ethylene glycol with urea under reduced pressure with heating and ZnO catalyst.

2)Filter solution hot to remove catalyst.

3)Cool solution in freezer and decant top unreacted layer.

4)Add 25% volume of xylene and distill to remove some of the ethylene glycol.

5)Change receiving flask and continue distillation to distill off the ethylene carbonate.

The azeotrope avoiding the close similarity in boiling point of ethylene glycol and ethylene carbonate and the freezing removing the bulk of the ethylene glycol before hand. Supposedly according to the patent the yield based on urea is 100% so there should be no carry over of this unreacted product.

BromicAcid - 10-1-2006 at 21:09

I have a question on the practical manufacture of methyl carbonate.

The usual industrial way is the reaction between methyl alcohol and carbonyl chloride as discussed before:

2MeOH + COCl2 ---> (MeO)2C=O + 2HCl

The question is, what are the reaction conditions? Is this a gas phase reaction, a liquid phase reaction, is there some solvent to mediate it? I mean, it makes 2 mol of HCl, I would think a strong acid like that would cleve the formed ester, maybe it has to be absolutely anhydrous?

Does anyone know anything about the practical use of this reaction?

The_Davster - 10-8-2007 at 22:37

I am bringing this thread up from the depths to ask the same question as Bromic. I have been searching for references to that reaction, with no luck at all. I am especially interested in unsymetric carbonates from phosgene(or triphosgene) All I have been able to find is that the chloroformate ester formation rate constant is huge compared to the next mole of alcohol reacting with the chloroformate, allowing the synthesis of unsymetrical carbonates.

(no this is not for home

...phosgene use would be insane at home)

S.C. Wack - 11-8-2007 at 05:00

One of my "65 selected chemical reviews" at FTP#2, rev_64_645_1964, has some refs to that date on a few specific chloroformates and their reactions with certain alcohols. It vaguely states that unsymmetrical (or symmetrical) carbonates would be made from one alcohol and phosgene at low temperature without solvent to give the chloroformate and HCl, then adding an alcohol and refluxing. It sounds quite easy, and DMC seems to be simply made by adding phosgene to methanol at STP and refluxing, though some refinements are mentioned.

One would have to hit Beilstein or CA for more modern (and in English) refs than the old German and French ones. Volunteers?

Sauron - 11-8-2007 at 05:51

See my thread on methyl chloroformate (chlorocarbonate.)

Ethylene carbonate is a very cheap commodity now. While it is noce to know how to make it from EG and urea, it is hardly worth the bother.

One of the more interesting things to do with ethylene carbonate is to prepare ethylene sulfide from the carbonate and KSCN. See Org.Syn. You need a better than usual pump for this (1 to 0.1 torr)

The product is the smallest cyclic sulfide. It can be stored at RT for several weeks without polymerizing.

[Edited on 11-8-2007 by Sauron]

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BromicAcid - 11-8-2007 at 08:03

The fact that ethylene carbonate is very cheap is irrelevant to this thread. It was fun to make and gave me an exercise in reduced pressure stripping of NH3.

Here is some more random information that I found out since my last post which is unrelated to the previous few posts but are in the spirit of the thread. I have read that some polycarbonates can be 'cracked' by refluxing with high temperature solvents and then extracting the reaction mixture with different solvent systems. Not much to go on but might be an interesting way to go considering the availability of polycarbonates.

And on the uses of carbonates front, I ran across this tidbit of information:

Quote:

A process is provided for producing oxalyl chloride by first photochemically chlorinating ethylene carbonate to form tetrachloroethylene carbonate and hydrogen chloride and then decomposing the tetrachloroethylene carbonate to oxalyl chloride and phosgene. The chlorination reaction is performed in a reaction vessel having an illuminated side arm conduit of narrow cross-sectional area. Chlorine is introduced at the base of the conduit and together with the hydrogen chloride evolved in the reaction continuously circulates the ethylene carbonate-chlorine reaction mixture through the narrow reaction zone within the conduit. The temperature of the ethylene carbonate-chlorine reaction mixture is controlled in the range of 70.degree. to 100.degree. C., preferably in the range of 79.degree. to 81.degree. C. The tetrachloroethylene carbonate is decomposed by heating same with a catalytic amount of a tertiary amine or amide.

From US Patent #4390708

That entry turned up in a book I was reading on phosgene and its derivatives (diphosgene, triphosgene, etc) and was somewhat praised for an easy prep of oxalyl chloride, though on the lab scale producing this much chlorine might be a pain in the butt, it does constitute a use for an organic carbonate.

Sauron - 11-8-2007 at 08:30

As has been mentioned, but which can never be emphasized enough: phosgene is an insidious poison that combined great volatility with life threatening effects at sub-irritating concentrations with symptons not appearing for long enough to allow a lethal exposure to take place while the chemist is blissfully unaware.

For that reason, all operations that require phosgene or generate phosgene (like this one) not only need to be done in a good hood and with a scrubber, but a phosgene detector/alarm should be employed (these are commercially available) and a SCBA worn.

Please note that surplus military gas masks are ALMOST ALWAYS out of use-by date. The AC filters are ineffective and the mask material likely compromised. The many Soviet masks in commerce are invariable worse than useless.

A SCBA uses compressed air not filters.

All in all, just about any alternative to phosgene is better. Diphosgene is still toxic but is much less volatile and is not insidious. It is however a pain to make. Triphosgene is a solid and even more of a pain to make, and is much less reactive than either diphosgene or phosgene.

Even if you are fully protected, have a care for your family, your neighbors. Murphy was an optimist, remember.

My personal opinion is that it is better to chlorinate methyl formate photochemically to diphosgene, rather than to generate phosgene to prepare methyl chloroformate and then chlorinate that. Managing the explosion risk of Cl2-methyl formate is IMO less dangerous than working with phosgene. Call me a coward.

Sauron - 11-8-2007 at 09:56

@BromicAcid, I did not mean that you were wasting your time. Have fun. It is an interesting preparation.