Sciencemadness Discussion Board

Deep Eutectic Solvents (OTC Ionic Liquid)

ShadowWarrior4444 - 15-5-2008 at 18:11

Deep Eutectic Solvents have had a bit of my attention for quite some time, the applications are quite intriguing. They are also exceedingly easy to make at home with cursory access to a nutrition store. The prime compositions of this type of solvent usually use Choline Chloride, though literature seems to indicate that any "clunky organic molecule" will work. (Quaternary amine chlorides are much preferred, usually paired with an a hydrogen donor to form the solvent.)

Notable properties:
Choline Chloride+Urea has an MP of 12C
CC+Glycerol at -35C
CC+Phenol at -40C

These solvents are electrically conductive, can dissolve cellulose, and many metal salts in large quantities. DES can also dissolve metal oxides and carbonates.

I personally have made and prefer the CC+Glycerol variety, and have recently been testing its stability in various environmental conditions. So far, the formulation has remained unfrozen and unchanged for 6 months at -15C. The current formulation I am using contains a bit of water due to the hygroscopic nature of CC, though dehydration of the CC before hand may make a very useful anhydrous solvent.

Due to its ability to dissolve a wide variety of salts, as well as its electrical conductivity, I have been considering it for use in an attempt to produce metallic sodium from electrolysis of a sodium salt in an unreactive anhydrous solvent as mentioned in:
http://members.aol.com/bromicacid/sodium/14.jpg
http://members.aol.com/bromicacid/sodium/15.jpg

A versatile solvent such as this may provide very unique opportunities to the home chemist.

(Note: I have also added the solvent to a collagen based gelatin before gelling. It has remained stable at 5C for 6 months as well.)

Useful reference:
http://www.rsc.org/Education/EiC/issues/2005_Jan/salty.asp

[Edited on 5-15-2008 by ShadowWarrior4444]

-jeffB - 15-5-2008 at 19:05

Me too! Me too!

I actually ordered a couple of Kg of USP choline chloride, mostly for this sort of experimentation. You can apparently deposit aluminum from salts dissolved in these mixes, and I wanted to see just how far along the electromotive axis I could go.

Unfortunately, choline chloride is EXTREMELY hygroscopic, in fact deliquescent. Leave some crystals scattered on an exposed surface and they'll quickly become droplets, even when the relative humidity is fairly low. I was easily able to make an aqueous solution containing 1g of choline chloride per ml. I think it's pretty darn close to "miscible in all proportions" with water.

The choline chloride I got was in the form of caked crystals. The papers I've seen talk about recrystallizing it from anhydrous ethanol, which I don't happen to have on hand, but may be able to get through my place of work.

I made the 2:1 molar eutectic with urea. It formed easily, but after sitting sealed at room temperature for a day or so (well above 15C), white waxy discs started to form at its surface -- it looked for all the world like it was starting to get moldy. Within a few days, most of the mix had solidified. I reheated it and everything went back into solution, but the phenomenon repeated upon cooling. Within a week or so, the entire mass had solidified.

I suspect I'm seeing this because there was too much water in the mix, and there's some component that freezes out of the ternary system. FWIW, my mix didn't do a darn thing to cellulose in the form of absorbent paper.

I intend to try the glycerol (and maybe ethylene glycol) mixes as well, but I think I'd better pursue choline chloride dehydration first.

12AX7 - 15-5-2008 at 19:09

What are the redox potentials for those solvents? I think you're going to have quite a problem bringing metallic sodium or elemental chlorine into contact with them. For instance, tetramethylamine hydroxide (or was it ethyl or butyl?) cannot be prepared pure, one of the alkyls breaks off. I forget what the reaction is called, but a cursory review of said chemical should find it.

Tim

ShadowWarrior4444 - 15-5-2008 at 19:37

Quote:
Originally posted by -jeffB
I made the 2:1 molar eutectic with urea. It formed easily, but after sitting sealed at room temperature for a day or so (well above 15C), white waxy discs started to form at its surface -- it looked for all the world like it was starting to get moldy. Within a few days, most of the mix had solidified. I reheated it and everything went back into solution, but the phenomenon repeated upon cooling. Within a week or so, the entire mass had solidified.


I had used a 1:1 ratio of glycerol to choline chloride, as is described in the literature, though 2:1 mixtures are also mentioned. It may be advantageous to attempt a series of mixtures, from 1:1 to 2:1 with urea, conducting a melting point test on each one, as well as stability tests. It should melt at 12C when the optimum ratio has been reached.
(Note: Found in journal article 3: "The DES of particular interest is a 2:1 molar mixture of Urea to Choline Chloride.")

Finding information on the redox potential of these ionic liquids hasn't been particularly easy, mainly papers that require purchase come up, and not all may be useful. Some I have come across:

Electroless deposition of metallic silver from a choline chloride-based ionic liquid: a study using acoustic impedance spectroscopy, SEM and atomic force microscopy
http://www.rsc.org/publishing/journals/CP/article.asp?doi=b7...

Application of ionic liquids to the electrodeposition of metals
http://www.rsc.org/publishing/journals/CP/article.asp?doi=b6...

148 U.S. Department of Energy Journal of Undergraduate Research
http://www.scied.science.doe.gov/SciEd/JUR_v7/pdfs/Chemistry...
Replicated here: http://www.scied.science.doe.gov/scied/Abstracts2006/chem.ht...

Ion transfer processes at ionic liquid based redox active drop deposited on an electrode surface
http://www.rsc.org/publishing/journals/CC/article.asp?doi=b5...

[Edited on 5-15-2008 by ShadowWarrior4444]

woelen - 15-5-2008 at 22:53

I am quite sure that certain not too reactive metals can be deposited from these solutions, but I have severe doubts on alkali metals, earth alkali metals and aluminium. E.g. glycerol is an alcohol and this reacts with those metals by formation of a glycerolide (sp?) anion and hydrogen:

2ROH + 2Na --> 2RO(-) + Na(+) + H2

-jeffB - 17-5-2008 at 05:44

Quote:
Originally posted by woelen
I am quite sure that certain not too reactive metals can be deposited from these solutions, but I have severe doubts on alkali metals, earth alkali metals and aluminium. E.g. glycerol is an alcohol and this reacts with those metals by formation of a glycerolide (sp?) anion and hydrogen:

2ROH + 2Na --> 2RO(-) + Na(+) + H2


I'll have to go back and refresh my memory from the papers I pulled. I think the aluminum deposition was done from the choline/urea mix, but I'm not positive. I was also reading a survey of ionic liquids, and I might be remembering one of the aluminum chloride-quaternary ammonium mixes.

Here's one abstract claiming that choline chloride/urea lets you electrodeposit zinc onto magnesium, which is challenging to say the least in an aqueous environment:

http://www.sciencedirect.com/science?_ob=ArticleURL&_udi...

What I'm remembering, fuzzily, is that the potential window for electrodeposition from these mixes is much wider than from aqueous solutions, thus broadening the range of metals that you can electrodeposit. I agree that it's unlikely we'd be able to plate an electrode with potassium, but I wouldn't be surprised to find lithium within reach. If the choline chloride/miscellaneous organic mixtures don't support it, the eutectic mixes with anhydrous ZnCl2 and the like might, and they're still well within amateur reach. The need to exclude moisture is beyond my technical skill at the moment, but it shouldn't be that much harder than running a challenging Grignard.

franklyn - 17-5-2008 at 22:48

Wikipedia cites a supplier which cites various reference papers here _
http://www.scionix.co.uk/index.php?option=com_content&vi...

Method for Dissolution of Metal - U.S. patent 5120523

Application of Ionic Liquids and Microwave
Activation in Selected Organic Reactions

http://herkules.oulu.fi/isbn9789514287190/isbn9789514287190....

.

[Edited on 18-5-2008 by franklyn]

vulture - 18-5-2008 at 00:42

I know someone who is doing research on ionic liquids for use in electrochemistry with rare earths. As these have quite high reduction potentials, the redox window of these materials tends to be fairly high.

panziandi - 18-5-2008 at 10:06

I remember reading a paper about side reactions of N,N'-dialkylimidazolium based ionic liquids undergoing reduction reactions under electrolysis conditions. Much like H+ forms H2 Im+ forms Im2 or something like that, I shall try to find a ref for it. That is what made me look into solvents which dissolve ionic compounds but are themselve not ionic. one class of solvent which screams out are the organic carbonates such as ethylene and propylene carbonates. I have several solvents to try and will trial out different anions to see which salts are most soluble also.

vulture - 18-5-2008 at 10:16

Diethylcarbonate is used in lithium batteries, so it should posses a very large electrochemical window.

ShadowWarrior4444 - 18-5-2008 at 13:44

Quote:
Originally posted by vulture
Diethylcarbonate is used in lithium batteries, so it should posses a very large electrochemical window.


Is Diethyl Carbonate produced in a method analogous to Dimethyl Carbonate? ('catalytic oxidative carbonylation of methanol with oxygen') [Pietro Tundo and Maurizio Selva (2002). "The Chemistry of Dimethyl Carbonate". Acc. Chem. Res. 35 (9): 706-16. doi:10.1021/ar010076f.]

Or must it be produced using Phosgene as detailed here: METHOD OF MANUFACTURING ALKYL CARBONATES

It also should be noted that the lithium battery solvent seems to be a chlorinated diethyl carbonate mixed with propylene carbonate in a range from 2-75%:
Chlorinated diethyl carbonate solvent for battery

[Edited on 5-18-2008 by ShadowWarrior4444]

panziandi - 18-5-2008 at 14:11

Alcoholysis of ethylene carbonate with ethanol will yield diethyl carbonate. Ethylene carbonate is commercially available but it can be made from ethylene glycol (ethane-1,2-diol) and urea. Although Ethylene carbonate is a great solvent for ionics too... There is mention of this in the forum elsewhere also google yields useful hits too... :)

franklyn - 10-7-2008 at 00:13

I am very much interested in using choline chloride solvation for anhydrous electrolysis,
but I am stymied in obtaining it. You would think that a nutritional supplement for livestock
and poultry should be easier to get than in 25 kg bags,
http://www.nutrimart.com/Bulk/Nutraceuticals.htm
and these are usually only 50 to 60 % content the balance being silica so you're buying
sand besides. The alternative is to buy it from a Chemical company but for what they
charge one could buy it in the big 25 kg bag.
http://www.coleparmer.com/catalog/productsearch.asp?search=c...
Another has it at more reasonble cost but won't deliver to a P.O.Box or residence.
http://www.usbweb.com/category.asp?cat=121&id=13410
The most common available form however is the bitartrate as the B4 vitamin, I suppose
treating with muriatic acid would produce the chloride but why go to the trouble since
it is already availble as a solution http://www.lef.org/newshop/items/item00541.html
http://www.amazon.com/Choline-Chloride-oz-Vitamin-Research/d...
containig 116 grams in a pint bottle. The problem then is how to dehydrate deliquescent
material into a dry powder. It melts with decomposition at just over 300 ºC so it may be
safely boiled to reduce the water volume then salting out by freezing the remainder.
CRC only indicates solubility in alchohols, but I found this added data online :
[color=darkgreen]- Freely soluble in water(1:4.5),slightly soluble in alcohol(1:90) and acetone. Soluble in
chloroform and insoluble in ether - [/color]So partitioning with cloroform or some other
chlorocarbon solvent may be a possibility. I am open to ideas or suggestions or if you
have a lead on a source send a private message.

.

-jeffB - 10-7-2008 at 10:47

Quote:
Originally posted by franklyn
The problem then is how to dehydrate deliquescent
material into a dry powder. It melts with decomposition at just over 300 ºC so it may be
safely boiled to reduce the water volume then salting out by freezing the remainder.
CRC only indicates solubility in alchohols, but I found this added data online :
[color=darkgreen]- Freely soluble in water(1:4.5),slightly soluble in alcohol(1:90) and acetone.


I'm not sure what temperature they had in mind for that 1:4.5 figure -- I think the real ratio is more like 1:0.45. I've made up some of mine as a nutritional supplement in a 1g/ml aqueous solution. Weigh out 50g (quickly, because it's gaining weight from atmospheric moisture while you work!), drizzle a few ml of water onto it, stir it while it "melts" into syrup, top off to 50ml. I generally have to microwave it up to maybe 40 C to get the last few crystals to dissolve, but once that's done it's stable at refrigerator temps (<5C).

I don't think it would be worthwhile to try to recover it from the LEF solution, and I think they preserve (contaminate) theirs with sodium benzoate anyhow. I used to get a 200g/l solution from Twinlab as a supplement, but I think that had pantothenic acid in the mix as well, and in any event they discontinued it years ago. Check your PM for the source I use now.

not_important - 10-7-2008 at 16:28

Quote:
Originally posted by franklyn
... The problem then is how to dehydrate deliquescent
material into a dry powder...


I'd try concentrating the water solution somewhat, then slowly dripping that into boiling toluene or xylene, distilling off the azeotrope. That should leave the salt as a powder. so long as you don't add too much water at any one time.

franklyn - 10-7-2008 at 22:17

Quote:
Originally posted by not_important
slowly dripping that into boiling toluene or xylene, distilling off the azeotrope. That should leave the salt as a powder.


That's a cute idea, but without a proper distilation setup, retort, condensor, etc.
inadequate ventillation can pose an explosion hazard. Another way would be to
freeze dry, sublimating the ice inside a sealed bottle with a good vacuum pump.
Also air drying by paint gun spraying into a large can circulating hot dry air.
I think that boiling to a viscous syrup and then blending with Epsom salt into a
doughlike paste and precipitating the hydrated epsom salt in methanol will likely
work.

.

ShadowWarrior4444 - 11-7-2008 at 13:32

Vacuum desiccator is likely the best option. Alternatively, if Choline Chloride doesn’t decompose at or below 100C, a wee bit of heat in a dry-air environment should do it. (Having alot of a strong desiccant in that environment would be prudent too.

Although, a bit of water may not be that troubling for many reactions--it may even aid the solvent for some as well.

FrankRizzo - 23-7-2008 at 17:06

Edit: Never mind..dumb question ;)

[Edited on 7/23/2008 by FrankRizzo]

franklyn - 5-11-2009 at 23:57

Some may find this article useful

Conductivities of AlCl3 Ionic Liquid Systems & Their Application in Electrodeposition of Aluminium
http://www.jproeng.com/qikan/manage/wenzhang/208139.pdf

.

franklyn - 10-2-2010 at 22:40

New organic solvents based on Carbohydrates
http://deposit.d-nb.de/cgi-bin/dokserv?idn=984940677&dok...

.

Attachment: Ionic liquids of hydrated salts patent WO0226381A2.pdf (1.4MB)
This file has been downloaded 1484 times

franklyn - 21-2-2010 at 00:19

Literature on ionic liquid solvents and low melting temperature salts.
All of the patents are conveniently packaged together in 2 attached zip files.


Low Melting Nitrate salt mixtures - US7588694.pdf
These low melting temperature nitrate salt mixtures are liquid below 95 ºC
Mol % of mixture
Li ,, Na , K , Ca ,
31 , 20 , 38, 12
25 , 16 , 47, 12
27 , 11 , 50, 12

Low Melting Sugar – Urea - Salt Mixtures as Solvents for Organic Reactions
Some are liquid at less than 70 ºC
http://www.rsc.org/suppdata/CC/b4/b414515a/b414515a.pdf

Physical Properties of Alcohol Based Deep Eutectic Solvents
Doctoral Thesis containing much data and references
https://lra.le.ac.uk/bitstream/2381/4560/1/2008harrisrcphd.p...

Ionic Liquid Electrochemical processing of Reactive Metals
Doctoral Thesis containing much data and references
http://circle.ubc.ca/bitstream/handle/2429/445/ubc_2008_spri...

Basic Ionic Liquids. A Short Review
http://www.ics-ir.org/jics/archive/v6/4/review/pdf/JICS-6-4-...

Res Metallica - Ionic Liquids: Solvents of the Future
Applications of Eutectic based Ionic Liquids

http://www.mtm.kuleuven.be/resmetallica/Abbott.pdf

Ionic liquids as solvents - US7183433.pdf

Method for producing Ionic Liquids - US20080251759.pdf

Separation of Liquid Eutectic by Crystallization on Cold Surface - US5814231

Ionic Liquids for separation of Azeotropes - US7435318

.

Attachment: US7435318 _ US20080251759.zip (1.4MB)
This file has been downloaded 1019 times

Attachment: USt5814231 _ US7183433 _ US7588694.zip (1.4MB)
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froot - 24-4-2012 at 23:22

Another interesting room temperature deep eutectic ionic liquid I found:

CaCl2-FeCl3

MrHomeScientist - 25-4-2012 at 07:12

That's interesting froot, but how exactly do you read that chart? Am I correct in interpreting that a mixture of ~67% FeCl<sub>3</sub> and ~33% CaCl<sub>2</sub> as solid salts will be liquid below 100 degrees C? I would think the hygroscopic nature of both salts would be problematic in such a mixture.

bquirky - 25-4-2012 at 09:45

Froot Have you tryed it out ?

Im curious how hot you have to get the mixture too kick-over into a euretic mix ?




froot - 25-4-2012 at 12:40

Nope I haven't tried it yet.
If I interpret the phaser diagram correctly at about 0.68 mole FeCl3 the mixture is liquid at about 12degC. The diagram is not straight forward so I may be overlooking something crucial to its properties.
In order to keep the mixture anhydrous I'd melt each constituent at relatively high temperatures before I mix them after which I'd decant the mix into an airtight container where it can cool.

AndersHoveland - 26-4-2012 at 01:06

So it appears the eutectic is 0.67 moles of FeCl3 to 0.33 moles of CaCl2, or on a weight basis, about 3 grams of ferric chloride for one gram calcium chloride
This is apparently a surprisingly simple ionic liquid.

CaCl2FeCl3.PNG - 61kB

[Edited on 26-4-2012 by AndersHoveland]

blogfast25 - 27-4-2012 at 14:10

Giver of bad advice, indeed :mad: :

Mole fraction FeCl3 = 0.67,

Thus 0.67 * 162.2 / (0.67 * 162.2 + 0.33 * 111 ) = 0.748 or multiply by 100 % = 74.8 weight %.

With 162.2 g/mol the molecular mass of FeCl3 and 111 g/mol the molecular mass of CaCl2.

White Yeti - 3-5-2012 at 14:57

I don't understand, aren't you just repeating what Anders said? Three grams out of four total represent iron chloride, ie 75% FeCl3 by mass?

barley81 - 3-5-2012 at 19:15

It is likely that Anders changed his comment when he was corrected by blogfast25.


[Edited on 4-5-2012 by barley81]

AndersHoveland - 4-5-2012 at 14:21

Quote: Originally posted by barley81  
It is likely that Anders changed his comment when he was corrected by blogfast25.

That would not have been possible. The post from me is marked
[Edited on 26-4-2012 by AndersHoveland], while the subsequent post from blogfast25 is marked "posted on 28-4-2012". It would not have been possible for me to alter the date of my post after I had altered it. Whenever a member alters their original post, the forum automatically adds an [Edited by ... on ... ] date stamp to the bottom of the post. These can be deleted by altering the post again, but this just adds another new date stamp to the altered post. Even if I had deleted the whole post and posted a new post, then faked the [Edited by ... on ... ] stamp at the bottom, the "originally posted on" date would have been later than the next post by blogfast25, which was posted on 28-4-2012. The top of my post reads: "posted on 26-4-2012 at 21:06", while the top of blogfast25's post reads "posted on 28-4-2012 at 10:10". It is not possible for a member to alter these dates. It is also not possible on this forum to modify a post after 24 hours of time has elapsed.

In any case, I did not correct my post in response to any outside criticism, not that it matters.

[Edited on 4-5-2012 by AndersHoveland]

dann2 - 4-5-2012 at 14:26


I hereby find the defendant not guilty.

Dann2

[Edited by Dann2 on 1 May] ;)

barley81 - 4-5-2012 at 15:24

Sorry Anders. I apologize for misconduct.

White Yeti - 4-5-2012 at 17:30

Quote: Originally posted by AndersHoveland  
It would not have been possible for me to alter the date of my post after I had altered it.


Then there's only one possible explanation...

Time travel:D:D:D JK

MrHomeScientist - 23-1-2014 at 14:55

As part of my long-running experiment to obtain neodymium metal from magnets, I've started investigations into deep eutectic solvents inspired by this thread and some additional searching. You can find my latest progress on this project over at the neodymium thread, here: http://www.sciencemadness.org/talk/viewthread.php?tid=14145&...

====================

The liquid I wanted to start with was choline chloride / urea at a 1:2 mole ratio. I purchased both chemicals from eBay - the choline chloride (CC) as fairly damp transparent crystals, and the urea in prill form. The CC has a rather strong, disagreeable smell to it that reminds me of triethylamine. I didn't bother drying it this time, but I put a portion in a dessicator for future use.

I ground 8.0g urea down to a powder in a mortar and pestle, and loaded this into a screw-top test tube. I then scooped 9.5g CC into the tube on top of the urea. Stoichiometrically, I only needed 9.3g but I wanted to add a little excess to help account for the water it's obviously absorbed.

This setup initially looked like this:
1.jpg - 251kB

I then took this outside and gently heated with a propane torch, moving the tube in and out of the flame to reduce hot spots and prevent any decomposition. In retrospect, I should have put the urea on top of the CC instead since it has a lower melting point. As it was, the urea melted first and I had to slosh it around a bit to dissolve the CC sitting on top of it. Once everything was liquid, I stopped heating.

After cooling, an interesting thing happened: two layers of crystallization!
2.jpg - 197kB

You can see it had solidified on the top and bottom, but left a small band of liquid in the middle. I suspected that this happened because I did not mix the liquid well enough, and while it formed enough of a mixture to be liquid while hot, it was not close enough to the eutectic point to stay liquid at room temperature. So the top layer would be CC-rich eutectic and the bottom layer would be urea-rich, while the middle was "just right."

I then took the test tube and heated again until molten. This time I used a glass rod to stir everything around and mix thoroughly. Upon cooling, it stayed liquid!

3.jpg - 146kB

The liquid is very viscous, similar to syrup. The little black bits floating around are from when I hit the CC with the torch a little too hard and burnt a small bit of it.

I look forward to experimenting with electrodeposition of metals with these types of solvents. It sounds like the CC / glycerol mixture may be more suitable for this, but that's what experimentation is all about. More to come!

blogfast25 - 24-1-2014 at 14:07

Very nice indeed.

I'd try to dissolve LiCl into it. Easier to obtain than anh. NdCl<sub>3</sub>, just as a starting point...

Help with ionic liquids

Azane - 27-1-2014 at 21:17

So I've recently become interested in ionic liquids (mainly for applications regarding water-sensitive solutes), and I have a few questions.
According to a video I saw on youtube, some of them can dissolve transition metal oxides, which are normally insoluble due to the electrophilic nature of transition metal ions. My first concern is how strongly basic the oxide ions of transition metal oxides would be in solution. I'm worried about that because, at the moment, the only ionic liquid I can make, with the materials to which I have access, is ethanaminium nitrate, which would probably react with the oxide ions to form water, ethanamine, and a metal nitrate, or so I assume, considering that ammonium ions can be deprotonated by hydroxide ions (more weakly basic than oxide ions) to form water and ammonia.
And if water-insoluble oxides are soluble in ionic liquids, would other normally insoluble/water sensitive ions behave similarly, assuming that the solvent ionic liquid does not feature portions that are susceptible to reduction/deprotonation? For instance, could sulfide, nitride, imide, amide, phosphide, or ethynediide (acetylide) dissolve without reacting? Furthermore, could alkali and alkaline earth metals dissolve in ionic liquids, just as they do in anhydrous ammonia?

I am also wondering if there are other ionic liquids I could make. The various cations I've focused on are those formed by dimethylamine, ethanamine, 2-aminoethanol, 1,2-diaminoethane, and the methylated derivatives thereof. (All in all, that totals out to 28 unique cations, including those derived from the non-methylated precursors, except for those of dimethylamine, which would be methanamine and ammonia). Of these, I think that 1,2-diaminoethane and its derivatives could have the greatest likelihood of forming ionic liquids with the proper anion.
The anions I've considered include nitrate, sulfate, disulfate, dithionate, 2-hydroxyethanoate (glycolate), 2-aminoethanoate (glycinate), and 2-aminoethanesulfonate (taurate). Of these, I think that 2-aminoethanesulfonate could most likely form ionic-liquids.
Confirm or repudiate whatever speculation of mine that you can, and give any additional information you think could help, please.
Thanks.

MrHomeScientist - 28-1-2014 at 11:22

Allow me to direct you to the ionic liquid (deep eutectic solvent) thread that I resurrected a few days ago, located just a few posts below this one: http://www.sciencemadness.org/talk/viewthread.php?tid=10529

The thread itself may not answer your questions directly, but there are numerous links posted that may offer better insight. In any case, that would be the appropriate thread to post these questions in rather than starting a new one.

I just started exploring these solvents, so I do not have any answers myself as of yet.

Azane - 28-1-2014 at 12:38

Cool, thanks man.

MrHomeScientist - 30-1-2014 at 16:18

An interesting development: After making this a week ago and letting it sit in my lab, lots of needle crystals have formed in the liquid.

1.jpg - 137kB

Perhaps I didn't get the ratio of the components quite right?

blogfast25 - 31-1-2014 at 06:30

Quote: Originally posted by MrHomeScientist  
An interesting development: After making this a week ago and letting it sit in my lab, lots of needle crystals have formed in the liquid.



Perhaps I didn't get the ratio of the components quite right?


If so, the clear liquid probably is the eutectic.

MrHomeScientist - 3-2-2014 at 09:05

I made another batch of the eutectic yesterday, this time with choline chloride dried for a week or so over calcium chloride. I used 8.0g urea and 9.3g dry(ish) CC. I also mixed the two reagents before heating, which made things go a bit more smoothly. They've melted together nicely, but left a lot of tiny white specks throughout the liquid. I can't tell if these are bubbles or undissolved reactants - the viscousity of the liquid may be trapping the tiny bubbles from moving, if that is what they are. I'll observe this batch and see if any needles separate out. I'm hopeful that since I dried the CC I can get it closer to the eutectic point.

gatosgr - 7-4-2015 at 08:13

I've experimented with deep eutectic solvents as well, I've made some new DES of my own , can you tell me who sells choline chloride on ebay I can't find any.

[Edited on 7-4-2015 by gatosgr]

MrHomeScientist - 7-4-2015 at 10:44

I bought it from user adpy-chemistry. He doesn't have any items for sale currently, and if there isn't anything else on eBay right now then I don't know what to tell you. Just keep checking every now and then and everything shows up eventually!

Now that it's been over a year, I suppose I should update on what happened to the liquid from my last post :P It ended up clearing nicely and has remained liquid this whole time. So it looks like drying the choline chloride and intimately mixing the reagents before melting them together is the way to go.

j_sum1 - 7-4-2015 at 19:31

I am interested in this. I first heard of ionic liquids 9months ago. An extremely large number of possible applications. I will post a list when I get back to civilisation.
I had assumed that this was going to be out of reach as a home chemist. Now you tell me that it enters the realm of the possible and I am intrigued.
J.

gatosgr - 8-4-2015 at 04:08

Actually the problem with water being in the DES is that it becomes part of the DES in stoichiometric values and it will not evaporate , there's a research about that. Well I follow this guideline for making a DES : the uglier the molecules are the most likely they are to form a DES. Theres also correlation between the hydrogen bond donors and the DES strength. I have a hard time sourcing materials are there any online shops that sell chemicals except ebay?

[Edited on 8-4-2015 by gatosgr]

deltaH - 8-4-2015 at 09:19

Choline chloride is sold by the tanker-load to large commercial poultry farms... the battery chicken variety. I got 25l of 70% solution for free by doing a bit of phoning around and then getting them to tap 'a little' for me from the pipes after a tanker delivery.

The solution can be boiled outside in a skillet on a gas stove to dry until it's really viscous. Cooling down then forms a mass of crystalline material that should be sealed while still hot to prevent it from absorbing atmospheric water. This cooking also drives off fishy smelling impurities (I assume trimethylamine and friends?) so that one is left with pretty decent material.

I've had LOTS of fun with this including making choline soap (reported somewhere here) and yes, prepared many DES's as well.

gatosgr - 8-4-2015 at 09:29

I think you mean 25ml right? How pure did it get with rectystallization?

[Edited on 8-4-2015 by gatosgr]

deltaH - 8-4-2015 at 23:23

No, it was 25 litres! I had to buy and bring my own large 25l container for it. I arranged it beforehand with the people importing the material, I told them I was trying to make a new kind of choline soap with it and they were happy to supply it for free. They said I could simply tap what's left in at the bottom of a large storage containers and piping.

I did not analyse it for purity, the original material was feed grade, so I assume that's pretty pure. It has traces of fishy smelling compounds, but these are volatile, so they came off during the boiling down. There was also a small amount of orange-brown turbidity which I attributed to rust from tank storage. This settled to the bottom of my container over a couple of weeks so that the top liquid was completely transparent and could be decanted (the settling was slow because the liquid's viscous).

The recrystallized, slightly tacky white crystals, were odourless. All-in-all, I was very pleased with it all and it gave excellent results in a wide range of experiments I did with it.

gatosgr - 9-4-2015 at 06:01

You have pretty generous poultry farmers in your region.:D:D
ChCl although very cheap is almost impossible to get a hold of...does anybody know a substitute?

[Edited on 9-4-2015 by gatosgr]

DeIonizedPlasma - 1-5-2015 at 18:21

I have been interested as well in deep eutectic solvents lately, and managed to devise a way that I believe made choline chloride. Because of the difficulty everyone seemed to be having obtaining choline chloride, I looked for another choline salt with an anion that would form as insoluble a precipitate as possible when combined with a halide salt. In the end I arrived at choline bitartrate, which is much easier to find as a supplement. I dissolved 0.2 mol of the bitartrate in 500mL of water (0.2 moles was chosen semi-arbitrarily, IIRC it was very close to a saturated solution in 500mL and I wanted to make a small batch as a test for the first time) and a saturated solution of 0.2 moles potassium chloride. Potassium bitartrate is nearly insoluble in water, making it perfect for the isolation of the choline chloride.

I mixed the two solutions and within a few seconds, small crystals began precipitating which at first I thought were bubbles due to their size (They looked just like a carbonated beverage, but instead of rising they fell). I stirred the solution and it immediately turned completely white and opaque, precipitating vigorously. By the end there was about 3/4 inch of precipitate on the bottom of the 1L beaker, which would seem to agree with the vast insolubility of potassium bitartrate. I filtered this and got a large amount of clear fluid which held the ChCl.

To verify that I had ChCl, I put a few mL on a watch glass and put that on a hot plate. As it boiled, a white flakey salt was deposited on the edge of the glass. By the time it had almost boiled completely off, the edges were beginning to brown and blacken which would seem to confirm an organic salt. Once completely dry, I left the watch glass out and waited to see if it would deliquesce. I waited about 20 minutes and did not see anything happening, so I left it out overnight. When I came back the next day it had clearly deliquesced, and there were a few tiny grains of what appeared to be a salt scattered in the liquid.

At this point I am planning on dehydrating the majority of the solution and running tests for purity to see if I have something I can work with. The Potassium bitartrate is still very wet, which I assume is due to the choline chloride being mixed in even after filtration and holding moisture in the mass. If I can manage to dry this I will measure it to get an approximate value of yield.

Does anyone have suggestions for dehydrating a deliquescent salt? I am currently looking into the acetone method but would be interested in seeing what else you guys can come up with.

gatosgr - 1-5-2015 at 23:35

That's good news although you can use thrimethylglycine instead of choline chloride which is a supplement and easy to find.

You are right about this method I found it here:

http://www.pharmacopeia.cn/v29240/usp29nf24s0_m17300.html

The page says you can dry it at 120 C for 2 hours.

[Edited on 2-5-2015 by gatosgr]

jock88 - 2-5-2015 at 13:18

You can use vacuum to dry salts (can't you, or am I mistaken?)

DeIonizedPlasma - 2-5-2015 at 18:00

Yes, vacuum should dry salts. I am more concerned about isolating my choline chloride from solution without decomposing it as it is deposited on the sides of the vessel. I may try putting it under vacuum and seeing if the water boils off as I have a 25 micron pump.

Alternatively, does putting a solvent such as acetone into a saturated solution of choline chloride force it to precipitate out if it is insoluble in the second solvent? I know that people reclaim choline chloride from a DES with urea using acetone, but I am unsure regarding miscibility mechanics if anything would happen.

gatosgr - 2-5-2015 at 23:58

The pure choline chloride may be recovered from this aqueous solution by distilling off the water in the solution, preferably under reduced pressure. The choline chloride which is recovered will have a purity of 98% or better and may be converted to the U. S. P. grade of choline chloride merely by one recrystallization from a solvent such as isopropanol, isobutanol, etc.

You can also use a dessicator bag https://www.youtube.com/watch?v=XJFfS_YbbYI but it will take some time.

DeIonizedPlasma - 3-5-2015 at 11:54

That sounds like the best option right now. I don't actually have a vacuum adapter for distillation though, as due to some quirks of glassware obtaining I only have a graham condenser and a 180 degree adapter so that the setup ends pointing down.

I will make sure to update when I finally make a DES, as I have plans for a sodium chloride cell that I think would be most entertaining. Unfortunately I will not be able to continue with the process for another week as school gets really crazy this time of year :(

gatosgr - 4-5-2015 at 05:23

You can use a dessicator bag https://www.youtube.com/watch?v=XJFfS_YbbYI

DeIonizedPlasma - 5-5-2015 at 12:19

The issue I have with dessicators is the large volume of water that I need removed from the choline chloride solution. I have about 600 mL of water that I need to remove and that would take a ridiculous amount of dessicant and time to accomplish. If you have done a large scale dessication in the past I would love to hear how you did it, but currently I have never heard of the process being used in a situation like this.

aga - 5-5-2015 at 12:41

Perhaps a 5kg bag of broken rice sold as pet food could be dried in a normal oven, and then put inside a bin-bag along with your 600ml of water ?

[Edited on 5-5-2015 by aga]

MrHomeScientist - 5-5-2015 at 12:56

Wiki (I know, not the best source) lists the decomposition temperature of choline chloride as 302 C. So why can't you just boil away the water? If you want to be super careful, boil away until crystals start to form, allow to cool and precipitate more crystals, filter off, then dry the damp crystals the rest of the way in a desiccator.

gatosgr - 5-5-2015 at 13:05

Well regardless of how ridiculous this sounds you can make a cheap vacuum pump from an injection maybe you can lower the boiling point a bit although I havent calculated the pressure needed.

https://www.avs.org/AVS/files/10/1043c4c6-597a-498f-a45d-7f1...

thank me later:D

[Edited on 5-5-2015 by gatosgr]

[Edited on 5-5-2015 by gatosgr]

DeIonizedPlasma - 5-5-2015 at 19:03

Oh, I have a vacuum pump, I just don't have any of the requisite glassware for a vacuum distillation. This is still very interesting, I'll definitely save that for future use. Do you know what the maximum pressure achievable is with this pump?

gatosgr - 6-5-2015 at 01:24

Measurements indicate that this pump is capab
le of producing a vacuum of less than 5%
of atmospheric pressure (~.8 psi, ~40 torr).
This displacement (~ 60 sccm) and ultimate
pressure (.05 atmosphere) is adequate for
many of the standard vacuum demonstrations
(for example, crushing cans and bottles, expa
nding a marshmallow, etc.). It takes ~12
strokes to collapse
a 500 ml water bottle.

Arun2642 - 5-7-2015 at 11:12

Quote: Originally posted by gatosgr  
That's good news although you can use thrimethylglycine instead of choline chloride which is a supplement and easy to find.

You are right about this method I found it here:

http://www.pharmacopeia.cn/v29240/usp29nf24s0_m17300.html

The page says you can dry it at 120 C for 2 hours.

[Edited on 2-5-2015 by gatosgr]


Where did you find that trimethylglycine could be used as an alternative to Choline Chloride, I just tried making a DES with trimethylglycine 1:2 urea and trimetylglycine 1:1 urea. Both crystallized well above room temperature. The method I used was to melt the urea in an erlyn-myer flask, add the trimetylglycine, and swirl and reheat until it all formed a homogeneous solution. Is there a different procedure I should follow?

nlegaux - 9-7-2015 at 19:06

I have been able to create a Type IV deep eutectic solvent by heating a 1:1 molar ratio of AlCl3 and Urea under oil (the paper recommends nitrogen, but I don't have access to any). It worked well as a supercapacitor electrolyte for me, but when it gets down near room temperature it gets quite viscous. I used this paper as a reference: https://www.sciencemadness.org/whisper/files.php?pid=404131&...

nlegaux

gatosgr - 3-12-2015 at 03:06

Quote: Originally posted by Arun2642  
Quote: Originally posted by gatosgr  
That's good news although you can use thrimethylglycine instead of choline chloride which is a supplement and easy to find.

You are right about this method I found it here:

http://www.pharmacopeia.cn/v29240/usp29nf24s0_m17300.html

The page says you can dry it at 120 C for 2 hours.

[Edited on 2-5-2015 by gatosgr]


Where did you find that trimethylglycine could be used as an alternative to Choline Chloride, I just tried making a DES with trimethylglycine 1:2 urea and trimetylglycine 1:1 urea. Both crystallized well above room temperature. The method I used was to melt the urea in an erlyn-myer flask, add the trimetylglycine, and swirl and reheat until it all formed a homogeneous solution. Is there a different procedure I should follow?


The procedure I follow is to mix the two powders and heat it in a glass tube stirring with a glass thermometer , I normally don't let it exceed 110 degrees because compounds get pyrolyzed at that temperature.
If you're getting crystals then the ratio isn't right try trimethylglycine with something else.

[Edited on 3-12-2015 by gatosgr]

gatosgr - 3-12-2015 at 03:08

delete this double post

[Edited on 3-12-2015 by gatosgr]

deltaH - 3-12-2015 at 11:55

I have boiled down choline chloride solutions in an open pan. Eventually, it gets very viscous and forms films on the surface, this is the end-point, because on cooling the whole lot crystallised into crystals.

DeIonizedPlasma - 4-12-2015 at 21:14

I am having trouble understanding what you mean, deltaH. Choline chloride decomposes fairly easily and I feel that this would easily decompose it into fishy trimethylamine. Also, if it is getting viscous but there is still water, when you ultimately cool it down there will still be water in your batch of crystals. I suppose you could try vacuum filtration of this to remove as much liquid water as possible from the precipitate...

gatosgr - 5-12-2015 at 01:58

Just order a bag of trimethylglycine seriously..

deltaH - 5-12-2015 at 03:41

Quote: Originally posted by DeIonizedPlasma  
I am having trouble understanding what you mean, deltaH. Choline chloride decomposes fairly easily and I feel that this would easily decompose it into fishy trimethylamine. Also, if it is getting viscous but there is still water, when you ultimately cool it down there will still be water in your batch of crystals. I suppose you could try vacuum filtration of this to remove as much liquid water as possible from the precipitate...


I used to think the same as you, but then I tried it...

The boiling does smell fishy because in the beginning, the trimethylamine that's initially in there (typical of commercial samples) boils off, but later it doesn't smell fishy and the freshly crystallised crystals don't smell fishy at all either. It's only upon aging that the characteristic odour develops again.

For all intent and purposes, the crystals are dry except for trace water. If you do it, you will see.

I think the viscous liquid is molten choline chloride. I didn't measure the temperature, but it was @#!%! hot at the end (gas burner on full with a large skillet pan over the flame). Luckily I had a hunch that it was dry and so turned off the gas. I was pleasantly surprised when cooling that everything solidified into a crystal mass.

If you want it 100% dry, just place those crystals in a desiccator or use vacuum. At least there will be very little water to have to remove at that stage. There might be some ethylene glycol contamination, maybe the crystals can be washed with a non-polar solvent to remove that. I didn't bother.

My point... you CAN simply just boil it down, so long as you know where to stop it. Anyway, if you stop it too early and it's too 'wet' when cooled, just turn up the heat and heat some more.

Be careful, don't leave it open for cooling, ChC is damn hygroscopic and it might liquefy again if you leave it open.

I used a thick-base pan when doing it with the idea that a larger surface area would be better for the end stage when it's evaporating through a film more than boiling. I don't know if that's critical, just saying...

[Edited on 5-12-2015 by deltaH]

DeIonizedPlasma - 6-12-2015 at 21:47

Interesting that you did this on a gas burner. I hope you are aware that trimethylamine is flammable? I will try this with about 50mL of solution on a small pan. Did you slowly turn up heat or put it directly over the high flames? I can imagine something like that sending boiling ChCl solution everywhere.

deltaH - 7-12-2015 at 11:29

The amount of trimethylamine that comes off is very little (less than a couple percent?), it's a minor contaminant (although it smells horrible)... and I'm still here ;)

[Edited on 7-12-2015 by deltaH]

MrHomeScientist - 8-12-2015 at 06:57

I'm going to have to try that some time;my choline chloride smells awful! I thought that was just a property of that compound.

deltaH - 8-12-2015 at 12:59

The problem is that in time it will start stinking again, one of the big problems with my choline soaps :mad: I could make fresh stuff that smelled fine, but after a couple of weeks...

DeIonizedPlasma - 8-12-2015 at 20:17

Quote: Originally posted by MrHomeScientist  
I'm going to have to try that some time;my choline chloride smells awful! I thought that was just a property of that compound.


No, sadly choline chloride just likes to undergo a Hoffman degradation into trimethylamine (wonderful fishy smell), water, and ethenol/acetaldehyde from what I can find. Mechanism picture attached, taken from Electrochemical decomposition of choline chloride based ionic liquid analogues. Given that it requires the hydroxide rather than the chloride, I think it may be happening due to trace presence of water. I couldn't find more mechanisms on the degradation, but I wonder if there isn't one that would occur in anhydrous ChCl. Perhaps someone should attempt to seal some choline chloride with negligible water content and leave it out for a while to test if it produces trimethylamine.
gF3IG1v.png - 19kB

[Edited on 9-12-2015 by DeIonizedPlasma]

Texium (zts16) - 8-12-2015 at 20:25

Quote: Originally posted by deltaH  
I used to think the same as you, but then I tried it...

The boiling does smell fishy because in the beginning, the trimethylamine that's initially in there (typical of commercial samples) boils off, but later it doesn't smell fishy and the freshly crystallised crystals don't smell fishy at all either. It's only upon aging that the characteristic odour develops again.

For all intent and purposes, the crystals are dry except for trace water. If you do it, you will see.

I think the viscous liquid is molten choline chloride. I didn't measure the temperature, but it was @#!%! hot at the end (gas burner on full with a large skillet pan over the flame). Luckily I had a hunch that it was dry and so turned off the gas. I was pleasantly surprised when cooling that everything solidified into a crystal mass.

My point... you CAN simply just boil it down, so long as you know where to stop it. Anyway, if you stop it too early and it's too 'wet' when cooled, just turn up the heat and heat some more.
Then how is it that when a few months ago I decided to try drying some damp reagent grade ChCl using medium heat on a hotplate at my school, it burnt black within a few minutes and made the entire classroom smell like a burning fish market? (to the despair of the AP chemistry class I was sharing the room with)

deltaH - 8-12-2015 at 21:19

Because yours was probably near the end point of what I made. I stopped mine just before it got very hot and presumably would start decomposing. When it cooled, the liquid solified into a slightly damp mass of white odourless crystals. If I had continued to heat this, it would likely have burnt too, I simply acted on a gut feeling not to heat further. Turns out at that point most of the water was gone.

Due the extreme hygroscopic nature of ChCl, I would think that what you see as 'damp' crystals, is in fact a very small amount of total water (or perhaps ethylene glycol). What I mean, let's say saturated ChCl is 80% solution (thumb-sucked) and let's say that your crystals are dampened by 5% saturated solution, so the water causing the apparent dampness is just 20%*5% = 1% overall!

Finally, this IS very stinky! I did my boiling outside and slowly so as to limit the stink. No doubt some decomposition occurs in the process, but the product was pure enough for my needs, white crystalline and odourless... at least for a while.

I made very nice choline soap with it, the only problem is that after a while, my soap bars went from being odourless to very fishy, which made me really sad because it worked very well otherwise, i.e. it made a beautifully mild and gentle soap.


deltaH - 8-12-2015 at 21:29

Quote: Originally posted by DeIonizedPlasma  
Quote: Originally posted by MrHomeScientist  
I'm going to have to try that some time;my choline chloride smells awful! I thought that was just a property of that compound.


No, sadly choline chloride just likes to undergo a Hoffman degradation into trimethylamine (wonderful fishy smell), water, and ethenol/acetaldehyde from what I can find. Mechanism picture attached, taken from Electrochemical decomposition of choline chloride based ionic liquid analogues. Given that it requires the hydroxide rather than the chloride, I think it may be happening due to trace presence of water. I couldn't find more mechanisms on the degradation, but I wonder if there isn't one that would occur in anhydrous ChCl. Perhaps someone should attempt to seal some choline chloride with negligible water content and leave it out for a while to test if it produces trimethylamine.


[Edited on 9-12-2015 by DeIonizedPlasma]


I noticed that the hydroxide degraded when preparing my soap. In an early experiment, I simply dissolved my NaOH directly into my 75% ChCl solution letting it get hot as well. This was bad, it seemed to degrade rapidly and the solution turned yellow-orange!

I had the idea that removing water might prevent this, so I proceeded to dry my ChCl solution by the method described before, then dissolving the ChCl into 1.5 mol. equivalents (off the top of my head, check the choline soap thread) 100% glycerine and only then added and blended the NaOH. This was better behaved and didn't yellow much, at least in the time it took to make the soap.

Unfortunately, the soap still formed trimethylamine in time when used, possibly because it gets wet. Water might be speeding up the process.

[Edited on 9-12-2015 by deltaH]

gatosgr - 3-7-2016 at 22:20

Trimethylglycine also decomposes to trimethylamine but not very much.

Bezaleel - 2-8-2016 at 07:27

To those interested in RTILs (Room Temperature Ionic Liquids), the following book has been published recently: "Ionic Liquid Properties - From Molten Salts to RTILs"
http://link.springer.com/book/10.1007%2F978-3-319-30313-0
The last chapter is over 70 pages, and is all about RTILs.
I don't have full access, but maybe some mad scientists do.

Loptr - 2-8-2016 at 10:08

Quote: Originally posted by Bezaleel  
To those interested in RTILs (Room Temperature Ionic Liquids), the following book has been published recently: "Ionic Liquid Properties - From Molten Salts to RTILs"
http://link.springer.com/book/10.1007%2F978-3-319-30313-0
The last chapter is over 70 pages, and is all about RTILs.
I don't have full access, but maybe some mad scientists do.




Attachment: Ionic Liquid Properties_ From Molten Salts to RTILs.pdf (3.5MB)
This file has been downloaded 2521 times


zck1214 - 8-9-2016 at 09:25

I can not receive the eutectic solvent when using AlCl3 (anhydrous) and urea, Who can tell me why?

MrHomeScientist - 8-9-2016 at 09:46

Magic elves?

Need more details from you. No one can give a meaningful answer with no information to go on. Tell us what exactly you did and where it went wrong.

zck1214 - 8-9-2016 at 09:58

Quote: Originally posted by MrHomeScientist  
Magic elves?

Need more details from you. No one can give a meaningful answer with no information to go on. Tell us what exactly you did and where it went wrong.


The mole ratio is 1:1, N2 protection and nothing happened at room temperature, so I increased to 120oC, but only white powder can be found. Thanks.

MrHomeScientist - 8-9-2016 at 11:48

I'm just speculating, but urea doesn't melt until 130 C so you may not have heated it enough. Also ensure your aluminum chloride is really anhydrous; a difficult task if not in a professional lab. A nitrogen blanket is a good idea.

When I made the urea / choline chloride liquid, I melted the whole contents of the test tube with a propane torch. No precise temperature control. Maybe the AlCl<sub>3</sub> version behaves differently though.

zck1214 - 8-9-2016 at 12:07

Quote: Originally posted by MrHomeScientist  
I'm just speculating, but urea doesn't melt until 130 C so you may not have heated it enough. Also ensure your aluminum chloride is really anhydrous; a difficult task if not in a professional lab. A nitrogen blanket is a good idea.

When I made the urea / choline chloride liquid, I melted the whole contents of the test tube with a propane torch. No precise temperature control. Maybe the AlCl<sub>3</sub> version behaves differently though.


Thank you. I have tried higher temperature than 130 C, the white smoke on the top of tube exhibits the sample evaporated. Maybe my AlCl3 has not high purity, I will try again after vacuum drying. Thanks.

Type II DES (Choline Chloride with MgCl2.6H2O)

jktan26 - 6-7-2017 at 06:49

Has anyone able to synthesize type II DES ? I had tried every molar ratio according to journals ranges from 1:1 to 2:1 (MgCl2.6H2O:Choline Chloride). It seems like the products i got are always in milkish color (as shown in the attached photo) even though they are in liquid form at room temperature. Any idea to resolve this? :(


20170706_120216.jpg - 1.5MB

yobbo II - 28-3-2018 at 14:46


An eutectic, melting at 310C, in the system sodium perchlorate-barium
perchlorate occurs at 43 mole per cent barium perchlorate.

Zinov'ev, A. A., Cludinova, L. I., and Smolina, L. P., Zhur. Neorg. Khim., 1, 1850
(1956).

stamasd - 16-6-2018 at 05:33

I have experimented a bit with DES, in particular type 4 in my search for a method for electrodepositing chromium from Cr3+ salts (and also cobalt, though I have a simple system for that). I attempted to make DES by melting together CoCl2 and CrCl3 respectively with urea in equimolar amounts. The resulting liquids are very viscous, almost unmanageably so. They also solidify in mass at room temperature after a few hours. While they were still liquid I attempted electrolysis in order to deposit the respective metals onto cathodes of copper and nickel, with anodes of Co and Cr respectively. Not much success, some thin deposition occurred but the deposits were thin, non-adherent and dull.

stamasd - 28-7-2018 at 05:56

I've been thinking a bit more of revisiting DES and chromium deposition since I last stumbled on this thread and posted above. I found a PhD thesis from U. Leicester https://lra.le.ac.uk/bitstream/2381/38111/1/2014AlbarzinjyAA... which investigated this exact matter. After reading through it, it looks that they had the best results with non-eutectic mixtures (non-eutectic eutectics, is that a thing?) and additives.

The best results were with a 2:1 molar mix of urea:CrCl3.6H2O with addition of 20% w/w free water and 0.1M boric acid. The next best was with 2:1 molar mix of urea:KCr(SO4)2.12H2O with the same additives (20% w/w water and 0.1M boric acid). Deposition done on mild steel substrates at 40C for 1h, current density 150mA/cm^2 for the first and 130mA/cm^2 for the second. They are both reported to have much lower viscosity than the pure DES, better conductivity and produce bright chrome layers with good thickness and hardness.

I will try to reproduce that if I can. I don't have CrCl3, but I do have some chrome alum. They mention some other additives such as benzoquinone (not sure if 1,2 or 1,4) but they haven't been tested in that paper. Still searching for more data on that, may try later on.

Will report back with results when available.

stamasd - 29-7-2018 at 15:51

No results yet, but some pretty pictures with color changes.

I made the type 4 DES today on a 1/10 mol scale. Molar ratio 1:2 KCr(SO4)2:urea is 499g:120g, or at 1/10 scale 49.9g:12g
To this will add 20% water and boric acid to 0.1M. Per calculations that means 12.7g of a 2.5% boric acid solution that will bring both the required water and the boric acid.

Pictures below.



Chrome alum purple at the bottom, urea white on top.




1 minute in microwave, it started melting and changing color.




Another minute in microwave, it became really hot and it almost completely melted. Some alum crystals were stubborn. Nice emerald color. Consistency of syrup.




After cooling down a few minutes, in natural light. It became much thicker, like honey. Too bad I can't show the consistency in pictures.




After adding the boric acid solution and stirring for a few minutes. All dissolved. Same color. Much less viscous, similar to molten ice-cream.

I'll let it rest for a day or two (principally because I lack time to continue now) then onto electrolysis.





[Edited on 29-7-2018 by stamasd]

Loptr - 30-7-2018 at 14:45

For those ionic liquids that contain choline hydrochloride, is there a problem with the liquid absorbing water from the atmosphere?

stamasd - 30-7-2018 at 15:41

I don't know. So far I've only experimented with types of DES that don't use choline (type 4). In fact until today I didn't even have choline chloride... and then the mail came with my bag of choline. :)

Urea is also slightly hygroscopic but not deliquescent. With the type 4 using urea I have noticed however somewhat the opposite problem. I think that if exposed to air, some of the water that is brought in as crystallization water of the alum evaporates, and that is one of the factors that leads eventually to the solidification of the DES. I have noticed that a previous sample of a similar DES to the one above (that one was made with 1:1 alum to urea) after less than a day had become very viscous at the surface, almost making a "skin" at the top. And then it started becoming a solid mass from the top down. That's why I think there was evaporation at play.

[Edited on 30-7-2018 by stamasd]

BaFuxa - 24-10-2018 at 10:32

Note on Type IV DES using urea as hydrogen bond donor : the urea decomposes fairly quickly. Once I used a MgCl2-Urea DES and everything started smelling ammonia when I heated it up to about 150-200°C.



[Edited on 24-10-2018 by BaFuxa]

stamasd - 8-12-2018 at 10:25

Quote: Originally posted by BaFuxa  
Note on Type IV DES using urea as hydrogen bond donor : the urea decomposes fairly quickly. Once I used a MgCl2-Urea DES and everything started smelling ammonia when I heated it up to about 150-200°C.


FWIW there was no ammonia smell while making the DES above (urea-chrome alum) nor since. Perhaps it's a better idea not to heat urea-containing solutions until it decomposes. :)

stamasd - 27-1-2019 at 08:34

A little (belated) update on my chrome alum/urea/boric acid/water liquid above. I stuffed it in a drawer and it stayed there for a few months. I did get around yesterday and test it in an electrolysis experiment. Using a stainless steel anode, a copper cathode, electrode area about 1cm^2 each, 10ml electrolyte, electrode distance average 2cm, adjustable power supply.

First observation, it does conduct electricity but much less than a saline/water solution. To achieve a current 0.1A I had to crank the voltage to about 9V. That's an impedance of 90 ohms, vs 5-10 ohms for a similar saline cell.

Second, I did get chrome deposited but not great. The deposit is uneven, with areas of copper still showing through, and in places the deposit is heavy, dark gray and dull (the copper cathode was thoroughly cleaned beforehand). OTOH the deposition is fairly quick, in less than 1 minute I got the deposit as above and further electrolysis time (up to 10 minutes) didn't improve on it. Using a lower current of 30mA resulted in a similar deposit forming in about 5 minutes, on which extra time (up to 30min) didn't improve.

Third, the chromium plating isn't great. It is reasonably adherent overall, but with weak spots where it can easily get dislodged by light polishing. Since the appearance of the deposit off the bath was quite ugly I tried to improve on it by polishing with a chromium oxide paste. Now, Cr2O3 is a pretty harsh abrasive, granted. But light rubbing of the paste over the plated area with only light pressure with the tips of my fingers removed large areas of the deposit. What was left is pretty strong and adherent and was not easily removed with extended polishing; its appearance is also very good and shiny after the extra polishing. But it only covers about 50% of the plated area, the remaining is gone.

No pics, I did not want to get chromium oxide onto my lens. :(

Not a complete failure, but also not very encouraging.

(edit) I did try higher current densities up to 200mA/cm^2 but that required higher voltages and led to significant gas generation with a lot of bubbling especially at the cathode, that's why I lowered it to 100mA/cm^2 and below; the plating obtained with the higher currents was also significantly worse with lots of pitting.

[Edited on 27-1-2019 by stamasd]

stamasd - 28-1-2019 at 15:44

I forgot to mention in the post above that the liquid's impedance seemed to vary nonlinearly with the voltage. Actually I should say that the cell impedance varies nonlinearly (as I used physical electrodes with it, not perfect electrodes). When varying the voltage from 0V to 24V, initially no current passed up until 4V; at 5V the current was 0.01A; at 6V, 0.03A; at 9V, 0.1A; at 15V, 0.2A; at 24V, 0.3A.

Attaching a little quick-and-dirty chart.
Almost looks like the current profile for a diode.



Capture.JPG - 29kB

[Edited on 28-1-2019 by stamasd]

DraconicAcid - 28-1-2019 at 22:22

That makes sense as far as I understand it- there won't be any conductivity unless there's a redox reaction at both electrodes, and that won't happen unless the voltage applied is enough to force the reaction to go.

Cou - 22-4-2020 at 10:18

If you are into ester chemistry, aroma compounds, the DES formed from choline chloride and glycerol can extract fatty alcohols from impure esters. E.g. it can extract 1-nonanol from nonyl acetate, a mixture which would result from fischer esterification which doesn't go to completion.

https://www.sciencedirect.com/science/article/pii/S004040391...

I am trying this out right now, i just prepared 80 mL of DES from 1:2 molar equivalent of ChCl and glycerol. I have samples of nonyl acetate that are contaminated with the smell of 1-nonanol. Will post a lab report after I try purifying them, but I can't analyze purity b/c no access to university instruments during COVID-19, only report the smell (nonyl acetate is mushroom, 1-nonanol is citrus)