Sciencemadness Discussion Board

synthesis of Propylene carbonate

symboom - 16-3-2011 at 23:09

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

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[Edited on 22.10.13 by bfesser]

ScienceSquirrel - 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.

Making propylene carbonate

Upsilon - 6-10-2013 at 13:06

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?

bfesser - 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).

elementcollector1 - 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]

Upsilon - 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.

Crowfjord - 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?

elementcollector1 - 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?

sonogashira - 6-10-2013 at 23:25

http://link.springer.com/article/10.1007%2FBF00608791

Mesa - 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.

deltaH - 7-10-2013 at 03:33

Also works for glycerine BTW.

Upsilon - 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.

Oscilllator - 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?

deltaH - 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

Upsilon - 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.

Upsilon - 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.

elementcollector1 - 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.

Upsilon - 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?

elementcollector1 - 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.

Nicodem - 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.

deltaH - 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]

WGTR - 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.

halogen - 20-10-2013 at 14:21

That's amazing WGTR! You didn't happen to capture a photograph did you?

WGTR - 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.

KNO3-LiNO3.jpg - 61kB

watson.fawkes - 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.

deltaH - 21-10-2013 at 07:48

Quote:
Do you have any idea at all what the polarization layer around an electrode is?
Very much so, the electric double layer at the cathode of a brine is made up of Na+ ions, not H+, (with some chlorides trailing behind that layer) that's why it's hard for H+ to be reduced in neutral water at high rates without a platinum cathode, cause the Na+ layer is so concentrated!
Quote:
...and that they're not naively valid when an electric field is present
What exactly are you trying to say? For H2O <=> H+ + OH-, are you saying that the electric field changes the activation energy for this or that it changes the equilibrium position for this or something completely different, because I don't understand your point?
Quote:
have no interest in spending time educating you,
I see, yet again, how you're not trying to educate me (rolls eyes).
Quote:
above nonsense counts as a hazard to others.
Indeed, electrochemical mechanisms are VERY dangerous... perhaps we should have this thread moved to energetics.

[Edited on 21-10-2013 by deltaH]

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Let's try this from a different angle, from this wiki entry on glassy carbon, I quote the following:

"Glassy carbon electrode (GCE) in aqueous solutions is considered to be an inert electrode for hydronium ion reduction:[6]
H3O+(aq) + e- <=> H(aq) Eo = −2.10 V versus NHE at 25 °C
Comparable reaction on platinum:
H3O+(aq) + Pt(s) + e- <=> Pt:H(s) Eo = 0.000 V versus NHE at 25 °C
The difference of 2.1 V is attributed to the properties of platinum which stabilizes a covalent Pt-H bond.[6]"

Now these are electrode potentials measured under otherwise standard conditions and already you can see just how hard it becomes to reduce H3O+ without an electrocatalytic cathode because without a catalyst, the first step is the formation of nascent hydrogen that is not covalently bonded to anything. Granted, glassy carbon is a crap electrode, but I'm just proving a point how it can influence the reduction of hydronium ions and that this has nothing to do with fields and the like.

So you see, even without bringing in kinetic effects, you already have a problem with reducing H3O+ on an noncatalytic electrode in neutral brines.

Then operate at appreciable rates and you make all these problems exponentially worse as you start introducing rate limitations and your concentration of H3O+ is completely dwarfed by Na+!

[Edited on 21-10-2013 by deltaH]

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[Edited on 22.10.13 by bfesser]

Nicodem - 21-10-2013 at 10:10

Quote: Originally posted by deltaH  
So you see, even without bringing in kinetic effects, you already have a problem with reducing H3O+ on an noncatalytic electrode in neutral brines.
...
Then operate at appreciable rates and you make all these problems exponentially worse as you start introducing rate limitations and your concentration of H3O+ is completely dwarfed by Na+!

Finally, the 2H+(aq) <=> H2(g) Estd. = 0V is defined only when H+ concentration is 1M, platinum is used as cathode and rates are miniscule!!!

Oh please stop it with these long facepalm replies already. Why do you bother instead of just posting something relevant with a reference, or a review like this one?
You don't understand much about electrochemistry, neither do you bother to learn first, so stop pretending, misquoting wikipedia, misunderstanding the concept of overpotential and abusing the misnamed "electrocatalytic electrode" concept.
You don't even understand that it is water that is being reduced, not the already present H3O+ or OH- (at the magnitude of the electric field at the electrode surface, the self-dissociation of water is irrelevant). And if it was Na+ being reduced, then the minimum potential for water electrolysis would be heavily electrolyte dependent, which it obviously is not unless one of its ions succumbs to an electrochemical reaction (in which case the electrolyte is no good for water electrolysis anyway). The electrolyte is there to make the solution conductive, that's all.

BobD1001 - 21-10-2013 at 12:00

Well I just ordered 16oz of propylene glycol. I already have plenty of urea as well. Ill I need is the Zinc oxide. Does anyone have any novel approaches to convert elemental Zinc into zinc oxide. Ive already found a couple methods online, just wanting to see if anyone has a creative method.

deltaH - 21-10-2013 at 12:12

Quote:
Oh please stop it with these long facepalm replies already.
As this is voiced as a request, I must respectfully decline.
Quote:
Why do you bother instead of just posting something relevant with a reference, or a review like this one?
I don't see how posting such a broad treatise on electrolysis supports any specific arguments you have made. I on the other hand have posted a specific link with a specific quote that is referenced to support my argument that if the cathode is inert but conductive, reduction of H3O+ becomes very hard indeed and your argument that why would it reduce Na+ when it can 'simply' reduce H+ at 0V is simply not true under the conditions that brine electrolysis is typically carried out with.
Quote:
You don't understand much about electrochemistry,so stop pretending
Your attempts to encite and enflame me are not working.
Quote:
misquoting wikipedia
quoted word for word and referenced
Quote:
misunderstanding the concept of overpotential and abusing the misnamed "electrocatalytic electrode" concept.
it would help if you explained yourself here, otherwise it's just an opinion.
Quote:
You don't even understand that it is water that is being reduced
I have repeatedly stated that that is what happens overall, but that at high rates, the pathway by which is might proceed may not be by the reduction of hydronium ions under the conditions of brine electrolysis. YOU stated that 2H+ => H2 has a Estd of 0V and so why would Na+ be reducing. That's what started this whole discussed and my argument was that mechanistically, this pathway is very minor under the conditions at which this process runs, i.e. neutral pH and ordinary fairly inert cathodes.
Quote:
And if it was Na+ being reduced, then the minimum potential for water electrolysis would be heavily electrolyte dependent,
No, I clearly stated that this is only significant at high rates and high rates is not where the minimum potential is measured. It is exactly because the easy pathway is exhausted at high rates that the reaction also starts to occur on less favourable pathways. At minimum potential, the easiest pathway dominates and no Na+ is reduced.
Quote:
The electrolyte is there to make the solution conductive, that's all.
I disagree but I concede this requires an experiment, so I propose the following that I will try to do as soon as I can:

I happen to have some choline chloride solution, which to my mind will perform poorly at high rates as there is no way that a quaternary ammonium ion can mechanistically reduce at the cathode, so that would FORCE direct water reduction on an electrode that is not electrocatalytic and so be the least efficient.

I will use graphite electrodes and carry out with 9V to ensure high rates (plenty of overpotential) and measure current using say 1M Choline chloride. Then I will repeat the experiment with identical setup and 9V, but use 1M NaCl and 1M HCl and measure current for each of those. I hope to show that the current will be highest for 1M HCl and comparable to NaCl and smallest for ChCl.

If you're done flaming, feel free to comment on such an experiment's methodology/construct which would be helpful.

[Edited on 21-10-2013 by deltaH]

Mesa - 21-10-2013 at 12:28

Quote: Originally posted by BobD1001  
Well I just ordered 16oz of propylene glycol. I already have plenty of urea as well. Ill I need is the Zinc oxide. Does anyone have any novel approaches to convert elemental Zinc into zinc oxide. Ive already found a couple methods online, just wanting to see if anyone has a creative method.


Burn it.
Throw MnO2 at it(Zinc batteries work on this concept, the outer casing is made of fairly pure zinc. I pulled a dead battery apart today and the entire outer casing disintegrated into white flakes.)
Dissolve in HNO3, then heat the resultant salt to 125*C(more efficient than burning the base metal, careful of the fumes though.)

papaya - 21-10-2013 at 12:41

Quote: Originally posted by Nicodem  

You don't even understand that it is water that is being reduced, not the already present H3O+ or OH- (at the magnitude of the electric field at the electrode surface, the self-dissociation of water is irrelevant). And if it was Na+ being reduced, then the minimum potential for water electrolysis would be heavily electrolyte dependent, which it obviously is not unless one of its ions succumbs to an electrochemical reaction (in which case the electrolyte is no good for water electrolysis anyway). The electrolyte is there to make the solution conductive, that's all.


Not to defend any side, I just don't understand one thing here - "The electrolyte is there to make the solution conductive, that's all", but to continuously conduct DC current there must be redox reactions present at the electrodes, otherwise the cell will quickly polarize and stop to conduct (capacitor behavior, if DC is applied), isn't it? If the electrolyte is not involved in the processes on electrodes(so only water is electrolyzed), then it'll not affect solution's conductivity for the reason I mentioned above and thus the conductivity must appear the same as for pure water (I'm not speaking about conductivity measurements with special instrument, where high frequency AC is employed thus no polarization occurs and no electrolysis also, consider only the DC case). Also the fact that with mercury cathode one obtains Na amalgam by electrolysis of NaCL solution indicates that somehow Na+ can be reduced from solution.

deltaH - 21-10-2013 at 12:53

Quote:
Burn it.
I wouldn't advise that as you might not obtain a material with a very high specific surface area and activity. Your best best would be to follow a specific procedure that is known to work. Would probably entail something like dissolving your zinc in dilute hydrochloric acid, then precipitating it in one fast go with the addition of stoichiometric amounts of dilute sodium hydroxide. Finally, let it stand for a while, decant, filter and wash + repeat before drying at moderate heat.

[Edited on 21-10-2013 by deltaH]

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@papaya Interesting point, I haven't thought about it in that way. I must admit, this is also challenging my preconceptions and understanding of the role of electrolytes, which is why I am very curious to conduct an experiment exploring this.

I couldn't at first glance find a decent paper where electrolytes were compared on the basis I'm thinking of. Usually these investigations into electrolysis are always conducted using amazing electrodes because they want to maximise efficiency and in such a case, you just need your electrolyte to provide the maximum ion conductivity.

All the mechanism takes place catalytically on the electrodes in the case of platinum and such and so probably the electrolyte plays no role in such cases other than to conduct the ions to and from.

I want to know what happens when Jo Bob (you and me) does this using simple electrodes we can get a hold of. I think in that case the choice of electrolyte may have a significant effect.

[Edited on 21-10-2013 by deltaH]

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[Edited on 22.10.13 by bfesser]

Nicodem - 21-10-2013 at 13:21

Quote: Originally posted by papaya  
"The electrolyte is there to make the solution conductive, that's all", but to continuously conduct DC current there must be redox reactions present at the electrodes, otherwise the cell will quickly polarize and stop to conduct (capacitor behavior, if DC is applied), isn't it?

The electrolyte is there in order to prevent the polarization - without it the cell would be only a leaking capacitor. You must have forgotten that that during the electrolysis of water there is a net flow of current from one electrode to the other because anions form at the cathode (2H2O + 2e- => H2 + 2OH-) and cations at the anode (2H2O => O2 + 4H+ + 4e-). In other words, the redox is occurring at two physically separated places, the oxidation at one side, the reduction at other. The electrolyte serves as a bridge connecting the reaction by moving the charge.
Quote:
If the electrolyte is not involved in the processes on electrodes(so only water is electrolyzed), then it'll not affect solution's conductivity for the reason I mentioned above

No, the nature of the electrolyte highly affects the conductivity because its ions move trough the solution.
Quote:
Also the fact that with mercury cathode one obtains Na amalgam by electrolysis of NaCL solution indicates that somehow Na+ can be reduced from solution.

Mercury electrodes have a huge overpotential for hydrogen (the highest of all metals). Overpotential is highly current density dependent and given enough current, this can be raised enough for the sodium amalgam formation. Actually, at high enough brine concentrations and low sodium concentrations in the amalgam, the hydrogen overpotential is higher than the sodium overpotential and thus the formation of H2 can be almost completely suppressed. See the short explanation in Electrochemical Engineering (H. Wendt, G. Kreysa), page 294.

deltaH - 21-10-2013 at 14:00

Quote:
Mercury electrodes have a huge overpotential for hydrogen (the highest of all metals). Overpotential is highly current density dependent and given enough current, this can be raised enough for the sodium amalgam formation. Actually, at high enough brine concentrations and low sodium concentrations in the amalgam, the hydrogen overpotential is higher than the sodium overpotential and thus the formation of H2 can be almost completely suppressed. See the short explanation in Electrochemical Engineering (H. Wendt, G. Kreysa), page 294.
I quoted Wiki's electrode potential for glassy carbon, an inert electrode, as being -2.1V, for the reaction H3O+ <=> H(aq), I presume this was determined in a 1M H3O+ solution as is the norm and that ?this does matter?, the concentration dependency of which can be calculated using the Nernst equation.

By your own arguments, this starts coming close to Na+'s reduction potential of -2.71V in a similar way than the mercury amalgam does (though even higher). Surely when you factor in the 10-7 H3O+ concentration of neural brines, why would you not be coming close to the reduction of Na+?

[Edited on 21-10-2013 by deltaH]

papaya - 21-10-2013 at 14:14

Thanks for the answer, I believe you must be knowing what you say, but the following is little complex for me to understand (or imagine)
Quote: Originally posted by Nicodem  

The electrolyte is there in order to prevent the polarization - without it the cell would be only a leaking capacitor. You must have forgotten that that during the electrolysis of water there is a net flow of current from one electrode to the other because anions form at the cathode (2H2O + 2e- => H2 + 2OH-) and cations at the anode (2H2O => O2 + 4H+ + 4e-). In other words, the redox is occurring at two physically separated places, the oxidation at one side, the reduction at other. The electrolyte serves as a bridge connecting the reaction by moving the charge.


So let's look at this more from the point of the physics. Total current = charge transferred per time. Only charge carriers in the solution are H+ and OH- that form at anode/cathode in redox processes as you wrote above - I said ONLY, because for example Na+ and SO4 2- while they exist in solution and will move in the electric field they'll NOT give their charges to anyone else thus they CANNOT transfer charge from one electrode to the other (otherwise we deal with Na2SO4 electrolysis, which you deny). Now, one situation - you have pure water, another - electrolyte added, do you say that in the presence of electrolyte H+ and OH- will move FASTER to opposite electrodes (because I = q/t and "I" is higher with electrolytes, only "real" charge carriers being H+ and OH- ), or do you say more H+ and OH- will form at electrodes per time that will move with the same speed?
Note I'm not wanting to disprove what you say, just want to clarify things for myself, and electrochemistry interpretations usually are confusing to me, I try to hold it simple.

watson.fawkes - 21-10-2013 at 15:30

Quote: Originally posted by papaya  
Only charge carriers in the solution are H+ and OH- that form at anode/cathode in redox processes as you wrote above - I said ONLY, because for example Na+ and SO4 2- while they exist in solution and will move in the electric field they'll NOT give their charges to anyone else thus they CANNOT transfer charge from one electrode to the other
If it's charged and mobile, it will move under the action of the electric field and thus be a charge carrier. If it's not a net charge carrier, then it only acts as a transient charge carrier. As a rule, the more charge carriers, the higher the conductivity.

Current flow does not require the bulk mass flow of its charge carrier. In a copper wire, the charge carrier are electrons, and no one expects the copper to move for there to be a current. If a water molecule ionizes in the bulk electrolyte, it is not the case that the two component ions need to move individually to the electrodes to complete their half reactions. These ions will collide with other water molecules, combining and resplitting in such a way that the charge can move faster than the component atoms.

Mesa - 21-10-2013 at 19:54

As with papaya, the vast majority of this stuff is way over my head so forgive me if the next question seems a tad stupid, but given glass becomes conductive when molten, does the previous poster's argument indicate that sodium would be obtainable by simply heating glass with a blowtorch and applying DC current?


Edit: Theoretical Sodium borohydride from a borosilicate glass melt?(Wikipedia hints at powdered borosilicate glass forming NaBH4 when reacted with sodium hydride, so I assume it'd need to be done in an atmosphere of hydrogen.)

[Edited on 22-10-2013 by Mesa]

elementcollector1 - 21-10-2013 at 20:44

Quote: Originally posted by Mesa  
As with papaya, the vast majority of this stuff is way over my head so forgive me if the next question seems a tad stupid, but given glass becomes conductive when molten, does the previous poster's argument indicate that sodium would be obtainable by simply heating glass with a blowtorch and applying DC current?


Edit: Theoretical Sodium borohydride from a borosilicate glass melt?(Wikipedia hints at powdered borosilicate glass forming NaBH4 when reacted with sodium hydride, so I assume it'd need to be done in an atmosphere of hydrogen.)

[Edited on 22-10-2013 by Mesa]


For the latter procedure, one must obtain sodium hydride first.
For the former, a blowtorch is not sufficient to make large quantities of glass truly 'molten'. Besides, there are much easier ways to obtain Na (namely, the electrolysis of the molten hydroxide, which can be melted with a blowtorch but really probably shouldn't - see the thread in Technochemistry).

papaya - 22-10-2013 at 00:33

Quote: Originally posted by watson.fawkes  

Current flow does not require the bulk mass flow of its charge carrier. In a copper wire, the charge carrier are electrons, and no one expects the copper to move for there to be a current. If a water molecule ionizes in the bulk electrolyte, it is not the case that the two component ions need to move individually to the electrodes to complete their half reactions. These ions will collide with other water molecules, combining and resplitting in such a way that the charge can move faster than the component atoms.

Completely agreed here - they don't have to travel the whole way rather will split next water molecule to a similar species, but what is the electrolyte doing then? And what is transient charge carrier you mentioned - seems it can explain things when you open brackets, thanks.

<hr width="80%" />
Quote: Originally posted by Mesa  
As with papaya, the vast majority of this stuff is way over my head so forgive me if the next question seems a tad stupid, but given glass becomes conductive when molten, does the previous poster's argument indicate that sodium would be obtainable by simply heating glass with a blowtorch and applying DC current?


Edit: Theoretical Sodium borohydride from a borosilicate glass melt?(Wikipedia hints at powdered borosilicate glass forming NaBH4 when reacted with sodium hydride, so I assume it'd need to be done in an atmosphere of hydrogen.)

[Edited on 22-10-2013 by Mesa]


First question - I think it can work, and once glass becomes conductive the high current applied may be sufficient to keep the glass very hot and conductive. Watch this made in russia experiment https://www.youtube.com/watch?v=p0E9Pd4CUBc
I can even speculate that the glass even doesn't have to be in molten state - it's enough to become ion-conductive, because I've heard about one experiment, when they turn on a light bulb, wait it to become hot, the apply a secondary high voltage DC - negative to the filament and positive to the glass from outside (don't know details), so after some time bulb becomes sodium "plated" from inside(electrons bombard conducting glass from inside), has anybody heard about this?

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[Edited on 22.10.13 by bfesser]

deltaH - 22-10-2013 at 01:14

Beautiful papaya, it never occurred to me that one could do this!

The fascinating thing about such a setup is that I would guess that at the cathode you would be producing sodium metal which probably quickly reacts with air to form sodium oxide and dissolves back into the glass, at the anode you probably produce oxygen gas. Now does this mean you can construct a hypothetical oxygen generator with such a setup if you make a much bigger cell? Air(oxygen) in on the left, pure O2 out on the right?

Because on the cathode side you consuming oxygen from air (provide electrode is exposed enough off course) and on the anode you are releasing very pure oxygen?

Electrochemical oxygen purification :cool:

[Edited on 22-10-2013 by deltaH]

<hr width="80%" />
Quote:
...split next water molecule to a similar species...
What did you mean by this?

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[Edited on 22.10.13 by bfesser]

Mesa - 22-10-2013 at 01:52

Quote: Originally posted by elementcollector1  
For the latter procedure, one must obtain sodium hydride first.
For the former, a blowtorch is not sufficient to make large quantities of glass truly 'molten'. Besides, there are much easier ways to obtain Na (namely, the electrolysis of the molten hydroxide, which can be melted with a blowtorch but really probably shouldn't - see the thread in Technochemistry).

Yes I realise this, I just performed that experiment today(unsuccessfully, poor soldering job on electrode caused it to fall apart.) I was only asking to get a better idea of the scope of non-aqueous electrolysis.

As for the requirement of sodium hydride in my theoretical synthesis, wouldn't a flow of hydrogen over the sodium as its forming on the electrode give at least a quantitatable(I checked google, turns out this is a real word:o) yield for proof of concept?

Still all theoretical, I'm having enough trouble with the molten salt electrolysis as it is. And I find pyrometallurgy much easier and more accessable given how much more forgiving it is in terms of temperature ranges.

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[Edited on 22.10.13 by bfesser]

watson.fawkes - 22-10-2013 at 04:37

Quote: Originally posted by papaya  
but what is the electrolyte doing then?
It's ionizing and deionizing, or rather forming transient neutral species with short lifetimes. Remember that a liquid is a mess of thermal motion as well. All it takes is a geometric bias for a given ion to form a neutral species nearer to one electrode and to split apart nearer to the other for it to act as a current carrier.

An energetic difference between an ionic type of bonding and a covalent one is the activation energy of transitioning between bonding and non-bonding states. Ionic bonding has much lower activation energy, so you'd expect more transient species to form as a result of random collisions. (You could put hydrogen bonding along this continuum as well, with ionic in the middle.)

papaya - 22-10-2013 at 05:12

Aha, turns out it's some sort of ELECTRONIC conduction in the solution, where the electron exits from cathode reducing water to OH- and from the latter it sometimes springs onto other molecule/ion and this way travels until gets the anode- more ions, higher probability to catch a "bus" to the next station! Because I can't call it something else than electronic conduction what is described (don't get angry if I misinterpreted something), prove me wrong.

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[Edited on 22.10.13 by bfesser]

deltaH - 22-10-2013 at 05:41

@Papaya

In nicodem's general review link he posted, there's this graph:



Notice that the current is very small until you hit the minimum potential for electrolysis to kick off... after that the current increases proportionally, so I don't think 'electronic conduction' occurs significantly in the electrolyte alone. Possibly the activation energy for electrons hoping on and off ions in much larger than say in a metal where there is a 'sea' of loosely bound electrons.

[Edited on 22-10-2013 by deltaH]

Nicodem - 22-10-2013 at 06:43

Quote: Originally posted by papaya  
Aha, turns out it's some sort of ELECTRONIC conduction in the solution, where the electron exits from cathode reducing water to OH- and from the latter it sometimes springs onto other molecule/ion and this way travels until gets the anode- more ions, higher probability to catch a "bus" to the next station! Because I can't call it something else than electronic conduction what is described (don't get angry if I misinterpreted something), prove me wrong.

No, no and no. It is not the electrons that transmit the current (except in the electrode and its surface). Like I already said, the ions of the electrolyte do this - that's why they are needed.
There really is no need to imagine wild things. You can imagine it as a simple train of charged particles (anions and cathions) each traveling to the corresponding electrode in order to maintain the local neutrality. The locomotive that forces them to travel is the electric field. This way every increase in charge at the electrode due to a molecule of water being electrolyzed is compensated by the slight movement of the equally charged electrolyte ions toward the opposite electrode. The ionic movement is dictated by the neutrality demand. Without this, the electrodes would just polarize due to the first few H+ and OH- formed at the electrode. The charge would not be able to be transmitted trough the medium due to lack of carriers and the electrolysis would stop immediately - you would have a capacitor.

Here is good page with short explanations:
http://chemed.chem.purdue.edu/genchem/topicreview/bp/ch20/fa...

There are some animated tutorials on electrolysis around. Here is one, but it only poorly addresses the role of the electrolyte:
http://www.youtube.com/watch?v=4bdFKyVM8fk
Try searching for a better tutorial, there must be some around, though the other I found were pretty much pathetic or simplistic to a degree of unfactuality.

watson.fawkes - 23-10-2013 at 05:50

Quote: Originally posted by Nicodem  
It is not the electrons that transmit the current (except in the electrode and its surface). Like I already said, the ions of the electrolyte do this - that's why they are needed.
If this were true, then there would be zero steady state current when an electrolytic cell is run at significantly lower than threshold voltage, say, 10 mV. Every such cell, though, has a mixture of both ionic and electronic conduction. Which is dominant depends upon lots of conditions: electrolyte composition, electrical field strength, temperature. In a 10 mV potential cell, the majority carrier in the steady state is indeed electronic. Ions act as a minority carrier (and not a non-carrier) because of their thermal motion.

In an ideal electrolysis cell, the bulk electrolyte has zero resistance, all the potential drop occurs in the polarization layers, only the ionic additions to the solvent react at the electrodes, etc. In that situation you would indeed have zero electronic conduction. In a well-run practical electrolysis cell, you have ions as the majority carrier and electrons as a minority carrier. In a not-so-well run one, there's more current than ionic conduction can provide (rate limited by ion drift velocity), and you get electronic conduction and electrolysis of the solvent.

The solvent itself also does provide some amount of electronic conduction. The conductivity of very pure water is small, around 10^-7 S/m, but that of insulating oils is around 10^-14 S/m and lower. You might be able to use tritiated water introduced at one electrode to measure the ratio of electronic to ionic conduction in such a situation, but I don't know if such work has been done.

Edit: added the units to conductivity values

[Edited on 2013-10-23 by watson.fawkes]

Nicodem - 23-10-2013 at 06:37

Quote: Originally posted by watson.fawkes  
If this were true, then there would be zero steady state current when an electrolytic cell is run at significantly lower than threshold voltage, say, 10 mV. Every such cell, though, has a mixture of both ionic and electronic conduction.

You are correct. The conductance that obeys Ohm's law is expected to be electronic in nature. This background current is however independent of the electrolysis reaction (it does not participate) and is generally magnitudes lower than the electrolysis working current. I should have stated that I was referring to the current involved in the electrolysis (the current that is generated due to H+ and OH- formation at the electrode). But then again, I'm not so sure anymore that papaya was asking about the electrolysis or in general, in which case you explained it better.

Upsilon - 3-11-2013 at 07:24

How well does PC work as a solvent in comparison to water when dissolving halide salts? Being able to dissolve certain substances is one thing, but to what extent is another.

elementcollector1 - 5-6-2015 at 14:13

Just put a bunch of (hydrated) yellow CeCl3 into some pure propylene carbonate. So far, it does not appear to be dissolving, but time will tell. Would cooling things down make it dissolve more?

blogfast25 - 5-6-2015 at 14:29

Stirring. Mild heating. That should do it.

But water is the enemy of electroplating REs. And 6 mol of water per mol of CeCl3 is A LOT!

elementcollector1 - 6-6-2015 at 23:41

Interestingly, *something* dissolved. I originally started with bright yellow CeCl3 flakes. When these were added to the propylene carbonate, they were bleached white - which is supposedly the color cerium chloride is supposed to be. If that's the case, I'm assuming the yellow color is from CeOCl, which seems to dissolve rather easily into the propylene carbonate (forming a greenish-yellow solution). The solution did not appear to be electrically conductive, but that may be because the amounts of CeOCl responsible for the yellow color were too low to establish any kind of actual current. I'll see if I can hydrolyze some of the white, 'pure' stuff tomorrow, and see if that CeOCl dissolves as well.

This is mainly a test to see whether the RE salts will dissolve to any appreciable extent in propylene carbonate, before I move on to trying to electrowin them. Therefore, water quantity isn't that important - for now.

elementcollector1 - 17-12-2015 at 17:14

Update on the above: The white stuff appears to be quite flocculent, and has broken up into small grains which do not stick together anymore. The supernatant liquid is a deep golden yellow in coloration, unchanged from when I last left it.

I get the feeling that the yellow color is not worth trying to isolate cerium from at this stage - if so, how do I purify the propylene carbonate for reuse?

I have a vacuum pump that can reach 5 Pa (apparently), but I don't know what the BP of propylene carbonate is at that pressure, and I'm afraid it might simply not condense under such an intense vacuum. I've heard of 'bleeder' pieces being introduced into distillation setups to raise the pressure a bit, but I don't know how to do this. Any advice?

halogen - 18-12-2015 at 12:45

solvents? if propylene carbonate is miscible with benzene (i checked it is) then water may separated phases salt in one, carbonate/benzene in the other. Unless you'll find you have a complicated mix, but how hydrophilic can propylene carbonate be?

anyway if you get it cold enough it will condense, look at pluto

[Edited on 18-12-2015 by halogen]

ah, yes, You see? miscible with benzene, ethyl acetate, etc., partly miscible with water. not micible with hexane. that's interesting.

https://books.google.com/books?id=oJy5wdzi0yUC&pg=PA178


[Edited on 18-12-2015 by halogen]

elementcollector1 - 18-12-2015 at 20:03

According to a nomograph, the BP of propylene carbonate at 5 Pa should be somewhere around 30.5 C. If that's the case, ice water should condense it, yes? Not sure how great the difference in temperatures needs to be to prevent too much vapor entering my vacuum pump.

Benzene is something else to try, if I can get my hands on it. For now, I'll stick with the resources I have.