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Author: Subject: synthesis of Propylene carbonate
symboom
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[*] posted on 16-3-2011 at 23:09
synthesis of Propylene carbonate


i was wondering if any one else has synthesized this compound C4H6O3 it is expensive to buy partly due to hazmat using the more cheaper available compounds synthesizing from urea and propylene glycol over zinc-iron double oxide catalyst.

[Edited on 17-3-2011 by symboom]

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Polar aprotic solvents


polar solvent that, unlike water, does not donate protons to the dissolved substances. It is relatively inert but has a high molecular dipole moment being helping to able to separate compounds in an inert environment

dimethoxyethane

Ethylene carbonate


not sure if electrolysis of sodium chloride will yield any sodium metal in these more common chemicals

acetone
ethyl acetate

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and maybe methyl ethyl ketone might also be more inert than acetone

<!-- bfesser_edit_tag -->[<a href="u2u.php?action=send&username=bfesser">bfesser</a>: merged 3 sequential posts]

[Edited on 22.10.13 by bfesser]
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[*] posted on 17-3-2011 at 12:22


Sodium metal is going to rip the shit out of acetone, MEK and ethyl acetate faster than you can say knife and sodium chloride is not more than trace soluble in any of them.
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biggrin.gif posted on 6-10-2013 at 13:06
Making propylene carbonate


I'm interested in making pure samples of alkali metals like lithium, sodium, etc. Doing some research I have discovered the wonders of propylene carbonate which, when a salt is dissolves in it and electrolyzed, will cause separation of the salt. The difficult part, however, is getting some of the stuff. It is not readily available to the average consumer. However, I have read that you can make it with propylene glycol, urea, and a catalyst, which are all cheap and easily obtainable. The details I have found regarding this procedure, however, have been ambiguous and unclear. Does anyone know more about this reaction that could help me out?
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[*] posted on 6-10-2013 at 15:21


Quote: Originally posted by Upsilon  
<u>I have read</u> that you can make it with propylene glycol, urea, and a catalyst, which are all cheap and easily obtainable. <u>The details I have found</u> regarding this procedure, however, have been ambiguous and unclear.
Provide the reference(s).



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[*] posted on 6-10-2013 at 16:55


Quote: Originally posted by bfesser  
Provide the reference(s).

Perhaps not the one OP read, but it's the same thing, really: http://www.sciencedirect.com/science/article/pii/S0167299104...
I'm interested in this as well, as this could be a better amateur route to sodium and potassium if one has a distillation/reflux setup. However, I've yet to find any mass amounts or stoichiometry... A balanced equation suggests CO(NH2)2 + C3H8O2 -> 2 NH3 + C4H6O3, but that was based off of this publication (http://ir.sxicc.ac.cn/bitstream/0/1501/1/573-576.pdf) which suggests that ammonia is the side product, and that a ratio of 1.5:1 PG to urea with about 1.2-1.3g of ZnO catalyst was optimal. It also claims an optimal yield of 99% propylene carbonate with catalysts of ZnO, Al2O3 or MgO.
So, if I had to guess how to go about it, it would be a one-pot synthesis with 1.5 mol of propylene glycol (distilled from antifreeze?), 1 mol urea (non-nitrate cold packs, should be cropping up everywhere), and 1.2g ZnO, in a 2/3-neck 250mL flask with a reflux condenser, thermometer, and a method of stirring (3-neck flask for overhead?). The reaction would be carried out at 170 degrees Celsius for 1.5 hours, and yield 102.09g of propylene carbonate, or about 85.1 mL.

This has been your daily spoonfeed.

[Edited on 10-7-2013 by elementcollector1]




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[*] posted on 6-10-2013 at 17:25


I don't even quite remember what the sources were. I know that some were discussions on this forum. Also, I don't mind at all buying the propylene glycol, urea, and zinc oxide straight-up. 4 liters of propylene glycol is roughly $25, 2kg urea is about $10, and 1kg of ZnO is about $13.

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[*] posted on 6-10-2013 at 18:02


There is this also, for ethylene carbonate. I am interested in this because it provides an amateur-friendly route to dimethyl carbonate. Just to clarify, are we talking about 1,2-propanediol or 1,3-propanediol?
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[*] posted on 6-10-2013 at 19:55


Quote: Originally posted by Crowfjord  
Just to clarify, are we talking about 1,2-propanediol or 1,3-propanediol?

If you mean the glycol, it's 1,2 -ethanediol.
That source is interesting, but is the vacuum really necessary?




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[*] posted on 6-10-2013 at 23:25


http://link.springer.com/article/10.1007%2FBF00608791
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[*] posted on 7-10-2013 at 00:05


This is something I've also been interested in for quite a while. I've recently bought 1kg of technical grade PG for specifically this reason.

I'd be interested to find out if the PG/Urea route can be done via microwave. I'll probably be testing it out in the next few days.
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[*] posted on 7-10-2013 at 03:33


Also works for glycerine BTW.



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[*] posted on 7-10-2013 at 12:55


So which method would be best for a home environment? Apparently synthesis of ethylene carbonate requires a vacuum, and I have not heard of this requirement for making propylene carbonate. If anyone can get this to work, I would be grateful if you posted the details of your endeavors.
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[*] posted on 7-10-2013 at 22:27


deltaH that results in the formation of glycerol carbonate, which is the same as propylene carbonate but has a hydroxide group. Am I correct in saying that this hydroxide group makes it a protic solvent, and therefore unsuitable because it will react with the metallic sodium/alkali metal?



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[*] posted on 8-10-2013 at 02:12


Indeed you are correct, not much use in terms of solvents in the presence of alkali metals.

However, the glycerine carbonate form very easily (see ref above).

In fact the only reason I discovered this was because I once microwaved a solution of glycerine and urea (for other reasons, amazingly not even related to my electrolysis work) and discovered a faint whiff of ammonia when hot. I suspected what was happening and indeed a quick google search turned up that patent. As you can see, the reaction occurs under fairly mild conditions in the presence of the catalyst (in my case slightly even without).

While not so interesting for making sodium or potassium, it is of potential interest in making liquid neutral surfactants at home with a carbonate head on it... they're polar as hell you know and you have the hydroxyl 'tail' to do a transesterification reaction with vegetable oils and they're 'green' as well.

I like to play around with, shall we say 'exotic' soap making lol





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[*] posted on 13-10-2013 at 17:22


Has anybody tried it out yet? I am excited to see some results. I would like to know what I'm getting in to before I spend $50 on ingredients.
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[*] posted on 14-10-2013 at 11:08


I just realized a crucial fact - any solution of a strong electrolyte will dissociate into its components when electrolyzed. For instance, when brine is electrolyzed, it produces free sodium and free chlorine. The reason that the sodium does not appear is because it instantaneously reacts with the water, explaining the formation of NaOH and hydrogen. So basically, any substance that will dissolve a salt and conduct electricity will break a salt into its elements. The issue is finding one that will not react with the elements formed.
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[*] posted on 14-10-2013 at 12:01


Quote: Originally posted by Upsilon  
I just realized a crucial fact - any solution of a strong electrolyte will dissociate into its components when electrolyzed. For instance, when brine is electrolyzed, it produces free sodium and free chlorine. The reason that the sodium does not appear is because it instantaneously reacts with the water, explaining the formation of NaOH and hydrogen. So basically, any substance that will dissolve a salt and conduct electricity will break a salt into its elements. The issue is finding one that will not react with the elements formed.

And it is quite an issue indeed - to find a solvent that will disassociate ionic compounds, but not be reactive. That's part of what makes propylene carbonate so valuable to the amateur.
I wonder how far this 'unreactivity' goes? If sodium and potassium are viable, that of course opens the way for lithium, calcium, strontium, barium, and magnesium (assuming the same conditions apply). But what about rubidium or cesium? Are they reactive enough to overcome this barrier? My imagining would be if one could electrolyze cesium out of this solution, one could place a container under the cathode to 'catch' the cesium that drips off (assuming that the solution is heated very gently to reach Cs's MP). Then, one could evaporate off the solvent under vacuum, and have a container full of Cs ready to use. Or, one could drip the cesium straight into an ampoule. Of course, this is armchair chemistry - I'd have to isolate some propylene carbonate and determine whether it reacts with Cs or Rb to any noticeable extent. If it doesn't, this would be an extremely viable route to the 'emperor of metals' given the OTC reagents needed.




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[*] posted on 17-10-2013 at 17:38


As far as the synthesis of the PC goes, do I just add urea to PG with some ZnO catalyst, while constantly stirring and heating?
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[*] posted on 18-10-2013 at 06:39


Quote: Originally posted by Upsilon  
As far as the synthesis of the PC goes, do I just add urea to PG with some ZnO catalyst, while constantly stirring and heating?

See one of my posts earlier up - it gives a very detailed account of what to do.




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[*] posted on 18-10-2013 at 07:47


Quote: Originally posted by Upsilon  
For instance, when brine is electrolyzed, it produces free sodium and free chlorine. The reason that the sodium does not appear is because it instantaneously reacts with the water, explaining the formation of NaOH and hydrogen.

That is pure nonsense (relative standard potential required to reduce Na+ is -2.71 V vs. the 0 V for reducing H+!). Why would sodium cations get reduced at the cathode when water is so much easier to reduce? It is just like saying that chlorine can oxidise fluoride anions. It can't.




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[*] posted on 18-10-2013 at 08:28


Quote:
That is pure nonsense (relative standard potential required to reduce Na+ is -2.71 V vs. the 0 V for reducing H+!). Why would sodium cations get reduced at the cathode when water is so much easier to reduce? It is just like saying that chlorine can oxidise fluoride anions. It can't.


Nicodem, if you use a platinum cathode, then yes, the water can be reduced at high rates because platinum can readily dissociate the water molecule on it surface to H* and OH* (*==surface species) with a low activation barrier. Then OH* + e- => OH- and desorbs, while H* + H* => H2(g) which also desorbs. This is because platinum is a catalyst for this process.

However, what happens with normal non-catalytic cathodes like graphite or steel? That's another story.

In neutral NaCl brine, the concentration of H+ is 10e-7M!!!

So there is simply not enough H+ for it to be reduced at the cathode.

The dissociation of H2O at appreciable rates without catalyst on graphite is very slow indeed, so I am afraid that I have to disagree with you and agree with the op that Na+, being in high concentration is the species which mechanistically is reduced when operating at practical and high current densities, to be immediately followed by reaction with water so that overall, it 'appears' as if 'just' the water was reduced, thermodynamically speaking this is all that matters as the internal pathway doesn't determine Estd., just the start and end states. At the very least, I might beleive that it's a combination of the two, where water reduction dominates at extremely low rates and gradually runs in parallel up to high rates where sodium reduction dominates. This is why the efficiency drops drastically as you up the rates (current density on the plates).

Na+ is neither created nor consumed in such a mechanism, so the Na+ reduction step is probably 'just' the rate limiting step. The point is that this is about kinetics and not thermodynamics.

I might be in error off course, but this makes more sense to me.

If, however, you were to use say HClO4 or H2SO4 as your electrolyte, then I would agree that H+ was being reduced because there is plenty of it, but the op was specifically referring to NaCl brines.

Finally, if you're operating at extremely low rates, then I might agreed with you that you give the water enough time to naturally dissociate on the cathode, but again... this concentration would be so extremely small that it's reduction would be merely academic and of no value at practical rates.

This is why standard potentials are measured at miniscule current, to minimise kinetic effects. Use of platinum in many cases also helps.

[Edited on 18-10-2013 by deltaH]

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On a more general note, there is a kinetic barrier for oxidising Cl- at the anode as well! Mechanistically, Cl- doesn't just magically become Cl2 in one step, it has to through a series of elementary reaction steps. Without a electrocatalytic anode, this can pose quite the activation energy barrier and so result in a large overpotential at high current density. This is why with pool electrolysers that act on brine, ruthenium dioxide coated anodes are preferred because they show the highest catalytic activity for chloride oxidation (lowest activation overpotentials) at high rates.

Many people think that dimensionally stable anodes like those are used just because of durability, but it's also about efficiency, especially because those electrolysers operate at high rates (current density).

[Edited on 18-10-2013 by deltaH]

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[Edited on 22.10.13 by bfesser]




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[*] posted on 20-10-2013 at 14:17


Even if lithium/sodium/potassium/etc reacts with propylene carbonate (or any other electrolyte/solvent), it may still be useful for isolating these alkali metals. For example, I've electrolyzed a eutectic mix of LiNO3 and KNO3 at 125-150C, and isolated small amounts of lithium that way. You'd think that lithium would react vigorously with nitrates, and it does. However, lithium is a solid below 180C. The oxide formed on the surface of the lithium is a fast ion conductor for Li+, and is insoluble in the nitrate melt. This protects the bulk of the lithium from the electrolyte. The problem that I had with this setup was that I didn't have a platinum anode. Pretty much anything else gets destroyed by the nitrates. Also, the molten salt has to be very anhydrous, or the oxide layer weakens.
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[*] posted on 20-10-2013 at 14:21


That's amazing WGTR! You didn't happen to capture a photograph did you?
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[*] posted on 20-10-2013 at 16:51


No, sorry. I still have the pea-sized piece of lithium that I made, but last time I checked it was heavily oxidized, even buried in petroleum jelly. I stored it under nitrogen, so I didn't expect it to last too long anyway.

However, I was able to conduct the experiment in a glass beaker on a hot plate, as the temperatures involved are not very high. The molten nitrate melt didn't attack borosilicate, and the electrodes were suspended freely in the melt, so they didn't touch the glass. The big problem, as I mentioned before, was the lack of a platinum anode. I was using stainless steel welding rod (mostly due to its low thermal conductivity), the erosion of which eventually turned the melt into a brown sludge.

This concept isn't anything new, though. If you look at how rechargeable lithium (metal) cells work, it's basically the same concept. There's that insoluble oxide layer that is there to protect the lithium from further oxidation (some cells use Li+ ion conductive glass). If you overheat one, the lithium melts, disturbs the oxide layer, and the cell goes into thermal runway (catching on fire).

I haven't tried this experiment with propylene carbonate, but as long as the lithium salt (maybe perchlorate) is soluble in it, and the oxide layer isn't, then it would probably work.

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[*] posted on 21-10-2013 at 07:09


Quote: Originally posted by deltaH  
In neutral NaCl brine, the concentration of H+ is 10e-7M!!!

So there is simply not enough H+ for it to be reduced at the cathode.
Do you have any idea at all what the polarization layer around an electrode is? Do you have any idea that the standard equilibrium constants are measured with an electric field of zero, and that they're not naively valid when an electric field is present?

I'm not particularly expert in electrochemistry, but at least I know the basics of the physical assumptions behind the mathematical model. I have no interest in spending time educating you, but the above nonsense counts as a hazard to others.
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