Sciencemadness Discussion Board - Salting out ethanol and distillation

Salting out ethanol and distillation

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deltaH - 8-3-2015 at 03:37

I happened upon this lab practical, see attached, with references.

In it is described how to salt out ethanol from water using potassium carbonate.

Their pic shows a pretty impressive separation.

I was thinking that this could be a very easy way to lower energy requirements and up the concentration and recovery of alcohol obtained, by the simple act of throwing in K2CO3 in the still. I would use much more than they did though, in fact, I'd throw in enough to make sure one saturates it at 100°C... that would be about 150g K2CO3/100g mash.

Afterwards, the residue can be dried out in the sun in wide open troughs to recycle the potassium carbonate for the next batch, easy peasy.

Also, instead of throwing salt into the still, another option is to use to reduce volumes prior to distilling by decanting and then distilling the organic phase only.

Why is this not done?

Reference:

http://ncsu.edu/project/chemistrydemos/Organic/SaltingOut.pd...

Attachment: Salting out ethanol.pdf (19kB)
This file has been downloaded 1359 times

[Edited on 8-3-2015 by deltaH]

j_sum1 - 8-3-2015 at 04:11

I like it. It seems very straightforward.
My first question is why potassium carbonate is used rather than sodium carbonate. The latter is so readily available.

But, yes I agree. It seems like a very simple and efficient procedure. There must be some reason why it is not commonly done. Probably only one way to find out.

deltaH - 8-3-2015 at 04:15

Thanks j_sum1.

As far as I can tell, potassium carbonate is used because of its much higher solubility, but you are right, in principle any salt that is very water soluble, inert(unreactive), non-toxic and cheap would do.

Wikipedia's solubility table is an easy way to look such things up for those that want to shake it up

chem_haruka - 8-3-2015 at 04:16

Yes,you can separation ethanol from water.
Separation of Alcohol–Water Mixtures Using Salts
www.ornl.gov/info/reports/1982/3445605458670.pdf

j_sum1 - 8-3-2015 at 04:35

Quote: Originally posted by chem_haruka

Yes,you can separation ethanol from water.
Separation of Alcohol–Water Mixtures Using Salts
www.ornl.gov/info/reports/1982/3445605458670.pdf

Ok. That's the firsttime I have seen a quaternary phase diagram presented as a tetrahedron. (page 20)

Praxichys - 8-3-2015 at 06:13

This method does not work well.

I tried this on a large scale and it was a bit of a disaster. Not only does it require huge amounts of carbonate to get any useful amount of ethanol, but the upper phase often has a significant fraction of water remaining, requiring the process to be repeated several times or followed by distillation.

I did this with a 1.75L bottle of 40% plain bottom-shelf vodka. Recovering the carbonate for reuse failed miserably. The water from the vodka contained traces of organics, enough to cause the carbonate to be contaminated by nasty-smelling char when I attempted to dehydrate it for reuse, ending up a golden brown color.

The first problem could be solved by saturating the aqueous layer with the carbonate, but this is impractical since K2CO3 solubility is 112g per 100g water at 20C. I would have needed something like 1200g of carbonate to separate 700ml of EtOH from the vodka.

Fractional distillation is the way to go, and maybe with a salt like K2CO3 in the still pot to bind the water ionically and improve separation.

deltaH - 8-3-2015 at 06:49

Thanks for the literature chem-haruka.

Thanks for the report Praxichys, some comments and questions:

Quote:

This method does not work well.
I tried this on a large scale and it was a bit of a disaster. Not only does it require huge amounts of carbonate to get any useful amount of ethanol, but the upper phase often has a significant fraction of water remaining, requiring the process to be repeated several times or followed by distillation.

It would be great to get some extra info here if you have it please. Specifically, what amount of K2CO3 did you use?

The 'followed by distillation' is strongly recommended because one would also want to reject the heads and tails that contain the impurities from naturally fermented alcohol, anyway.

The residual water is expected, but it MUST be significantly enriched in alcohol. Again, I'd throw the salt straight into the still pot in my principle version, as an aid to distillation.

Quote:

I did this with a 1.75L bottle of 40% plain bottom-shelf vodka. Recovering the carbonate for reuse failed miserably. The water from the vodka contained traces of organics, enough to cause the carbonate to be contaminated by nasty-smelling char when I attempted to dehydrate it for reuse, ending up a golden brown color.

I'd simply leave it in the sun to evaporate. Even on a large scale, this is easily done by digging a shallow trough, placing a large thick black plastic sheet to waterproof it, then pour in the tails, you could do hundreds of litres in this fashion if you wanted to, easily.

*This wouldn't work for mash tails though, only distillitate, though you could put it in a thumper though

Quote:

The first problem could be solved by saturating the aqueous layer with the carbonate, but this is impractical since K2CO3 solubility is 112g per 100g water at 20C. I would have needed something like 1200g of carbonate to separate 700ml of EtOH from the vodka.

Why? 25kg bags of K2CO3 from a chemical supplier is probably cheap and you can reuse it many times unless you use it in raw mash. If it's not easily available on a country-specific basis, we might be able to suggest substitutes that are.

Quote:

Fractional distillation is the way to go, and maybe with a salt like K2CO3 in the still pot to bind the water ionically and improve separation.

My thought here is that this could be the chemical version of a 'thumper'. Something that is really easy to do and can be done with existing still setups, but gives better recovery and concentration.

Perhaps the best way to use this is to put it IN the thumper, not the still, that way the mash residue doesn't foul it up as much and simply evaporation afterwards can be used to recycle the salt.

*****************************************************

Perusing the solubility tables... calcium chloride looks like a good alternative and it's available as de-icer AFAIK?

Also phosphate salts, e.g. monosodium phosphate?

[Edited on 8-3-2015 by deltaH]

Praxichys - 8-3-2015 at 08:28

Quote: Originally posted by deltaH

Thanks for the report Praxichys, some comments and questions:

Specifically, what amount of K2CO3 did you use?

I used 50% w/w with regard to the water in the ethanol.

Quote: Originally posted by deltaH

The 'followed by distillation' is strongly recommended because one would also want to reject the heads and tails that contain the impurities from naturally fermented alcohol, anyway.

Correct. My plan was to dry with 3A sieves afterward though, so I would then need an intermediate step to remove carbonate before sieving. I could save time and effort by not bothering with carbonate and going right to fractionating.

Quote: Originally posted by deltaH

The residual water is expected, but it MUST be significantly enriched in alcohol. Again, I'd throw the salt straight into the still pot in my principle version, as an aid to distillation.

Also correct, but keep in mind that heat will reduce the hold the K2CO3 has on the water. I have not tried this with carbonate, but this is generally why desiccants are strained before distillation. It will be better than no entrainer at all, but not perfect.

Quote: Originally posted by deltaH

Theoretically, yes. However, the reason the K2CO3 works, while other salts do not, is in part caused by the affinity the salt has for water. I do not think K2CO3 is deliquescent, but it must be close. You may have a hard time evaporating the solution without heat.

Quote: Originally posted by deltaH

25kg bags of K2CO3 from a chemical supplier is probably cheap and you can reuse it many times unless you use it in raw mash. If it's not easily available on a country-specific basis, we might be able to suggest substitutes that are.

True, but I got mine for $5/lb on eBay. It cost me about $10 to salt out 700ml of wet EtOH. For me, setting up a fractionating apparatus and taking the few hours required is free.

Quote: Originally posted by deltaH

Perusing the solubility tables... calcium chloride looks like a good alternative and it's available as de-icer AFAIK?
Also phosphate salts, e.g. monosodium phosphate?

CaCl2 will not work because both water and the smaller alcohols can form "hydrates" in the crystal structure. I am not sure about phosphates - it seems interesting to try! With monosodium phosphate you might end up with sodium diethyl phosphate though.

deltaH - 8-3-2015 at 09:15

Thanks Praxichys and good points. I'll try to think of better alternatives.

Once you get over the knee-jerk reaction, would nitrates be a definite no-no?

[Edited on 8-3-2015 by deltaH]

BromicAcid - 8-3-2015 at 09:42

Everything old is new again:

http://www.sciencemadness.org/talk/viewthread.php?tid=24299

deltaH - 8-3-2015 at 09:57

Oh crap, missed that, sorry guys.

Would anyone with glassware for distillation try a basic one stage distillation with some vodka and potassium carbonate and report on what's obtained as distillate?

Pretty please.

[Edited on 8-3-2015 by deltaH]

A small experiment: anh. zinc chloride + drinking gin

deltaH - 8-3-2015 at 10:55

Okay, performed a very simple experiment.

Hypothesis:

Saturating drinking gin with anhydrous zinc chloride will increase the alcohol concentration of the vapours enough to make it flammable.

Experimental:

I incrementally added anhydrous zinc chloride to a small amount of ordinary drinking gin (43% Vol.) until no more would dissolve. This took a lot of zinc chloride and led to a large increase of the solution's volume. No phase separation occurred at saturation. Then I poured some of this on a small ball of cotton wool and attempted to ignite it. Surprisingly, it wasn't flammable. :mad:

Discussion:

Zinc chloride is EXTREMELY soluble in water (432g/100ml, Wikipedia), but also in ethanol to a degree. The formation of the aqua complex from the hydration of anhydrous zinc chloride does consume a lot of water, but the resulting solution is diluted by the sheer volumes generated and the net result is that you don't have a large increase in the concentration of alcohol vapours, well not enough to ignite it, anyhow.

Conclusion:

When a salt is extremely soluble, it won't necessarily give you a great benefit of increasing the concentration of alcohol in the vapour at equilibrium.

*********

Now we already know that potassium carbonate works for bringing about a phase separation for ethanol-water, but I want to understand why. It looks like high solubility is not the only thing, it might also have to do with the interactions of the ions themselves, perhaps even their charge density. Understanding this better make allow us to select better salts for this purpose.

I'm thinking perhaps sodium orthosilicate (Na4SiO4) might work better, operating under the assumption that the higher the charge density on the anion, the better.

Problem is that sodium orthosilicate is not trivial to make... one would need to fuse silica gel and sodium hydroxide or carbonate at high temperature

[Edited on 8-3-2015 by deltaH]

Praxichys - 8-3-2015 at 11:24

It seems that my earlier failure could be mainly because I did not bring the water solution up to the saturation point with the K2CO3. As noted in the other thread, Magpie received back nearly all of the water after bringing the K2CO3 concentration to something near the saturation point.

Another good point was brought up in the thread regarding the cost of potassium carbonate. It could be made in-situ with super cheap sodium carbonate and KCl.

I will do some experimentation today.

deltaH - 8-3-2015 at 11:29

Quote: Originally posted by Praxichys

It could be made in-situ with super cheap sodium carbonate and KCl.

Yes I saw this, I also thought it was quite smart and the bonus is that the NaCl, not being very soluble, may precipitate, so not a worry there so much that it would have a negative effect.

Tripotassium phosphate could also be a good candidate going on the hypothesis that high charge density on the anion is beneficial. It very basic though, but I don't think that should be a problem. Unfortunately, my phosphoric acid is not at my current premises, but I can remedy that at a later date. I do have potassium hydroxide though, so making some K3PO4 should be easy peasy (once I get to my H3PO4 :mad:

)

Doing some reading, actually tripotassium phosphate looks like a very good candidate, it's solubility goes up a lot with increasing temperature and it's insoluble in alcohol (like potassium carbonate), plus it's a commercial fertiliser which means it probably can be bought in bulk bags on the cheap.

[Edited on 8-3-2015 by deltaH]

Zombie - 8-3-2015 at 12:05

Don't bother with any of those methods...

3A Zeolite is the easiest way to go, and you want to use it AFTER azeotrope has been reached, not before or during distillation. The cost is too high to fill a boiler.

To pre separate EtOH from water it is easier to freeze the water , an pour off the EtOH. Due to the volatility you really don't want to begin with more that a 40% ABV in a "still".

Controlled lab glassware is a different animal.

The Zeolite will pull out 100% of the water but exposure to air will replace it.

http://catalog.adcoa.net/viewitems/molecular-sieves/type-3a

[Edited on 3-8-2015 by Zombie]

deltaH - 8-3-2015 at 12:22

Zeolites are similar to using a desiccant, but it's an expensive way to absorb bulk water. It's great for removing that last little bit almost completely, but what I'm interested in is different. It's a salt you can use with low concentration alcohol that would result in the ability to distil a much higher concentration, simply, but not by desiccation, but by an ion effect reducing the water-ethanol interaction by favouring water-ion interaction.

As I said before, consider it as a 'chemical thumper', not a desiccant.

In fact, the way to use this would be to add the salt to a 'thumper' and not the still pot. In fact, probably even an empty 'thumper' with pure salt and nothing else.

I'll test out my hypothesis if using trisodium phosphate as soon as I can make it. That's also easily scalable because of the availability in bulk as commercial fertiliser (not the kind sold in garden centres, but the kind sold to big farma <<< see what I did there

)

[Edited on 8-3-2015 by deltaH]

Zombie - 8-3-2015 at 12:27

Quote:
"Zeolites are similar to using a desiccant, but it's an expensive way to absorb bulk water. It's great for removing that last little bit almost completely, but what I'm interested in is different. It's a salt you can use with low concentration alcohol that would result in the ability to distil a much higher concentration simply."

A properly designed, and run 6 plate 4" column will run Azeotrope in one run.
How much easier can it get?

Distilling is the same physical process no matter what the beginning concentration, and the same final product will always be azeotrope.

Other than a time machine... nothing will change that.

deltaH - 8-3-2015 at 12:33

Quote: Originally posted by Zombie

A properly designed, and run 6 plate 4" column will run Azeotrope in one run.
How much easier can it get?

I'm hoping a lot easier than that, equipment-wise. Also don't forget, there might be a significant energy saving as well... but that remains to be seen, of course.

deltaH - 8-3-2015 at 12:49

Speaking of energy, a thought just occurred to me. If one employed such a salt in a 'chemical thumper' then ideally you want the salt employed to give off heat on dissolving, else the liquid will just condense without vapours going off. Fortunately, tripotassium phosphate has a heat of fusion of +41.9 cal/g according to the CRC handbook, which is fantastic news. Potassium carbonate is +56.4 cal/g which means that should work even better from a heats point of view.

[Edited on 8-3-2015 by deltaH]

Zombie - 8-3-2015 at 12:53

You're beating a dead horse brother.

No matter what you do you need a reflux column.
Yes you can easily scale it down but no matter what you begin with the boil temp. is controlled by the EtOH mole fraction.

The more you raise the ABV, the more energy you need to release the vapor, and get the boiler up to that temp.

There's are couple thousand years of distilling to prove it out.

I suppose you could always chemically separate the water, and do away with distilling all together.

[Edited on 3-8-2015 by Zombie]

deltaH - 8-3-2015 at 13:02

Quote: Originally posted by Zombie

You're beating a dead horse brother.

Huh? The heat of vapourisation of ethanol is much lower than water, so the higher the ethanol concentration, the LESS energy it takes to vapourise.

Also, separation requires work, so the higher the concentration of the thing you want to enrich, the less work it's going to take to enrich it. To get 0.0001% alcohol solution to 50% is going to take a lot more work than getting a 40% alcohol solution to 50%.

I still think that forming a second phase where the alcohol is of higher concentration would improve distillation significantly, for a simple setup (still plus chemical thumper pot).

Shit, I really need to get away from using this term 'chemical thumper' very bad for PR. Perhaps a 'salt thumper' is better, in that case, might as well call it a Godzilla

[Edited on 8-3-2015 by deltaH]

Molecular Manipulations - 8-3-2015 at 14:11

I wonder what the the initial concentration if alcohol/water must be in order for salting out to work. I looked but couldn't find any information on the web.
I just tried it using potassium carbonate, calcium chloride and sodium phosphate (seperately) to salt what I believe to be about a 5% EtOH solution with water, sucrose and yeast (lots if contaminates I know) from a fermentation batch I've had going for a week or so. I just wanted to see if there's a way to raise the concentration before distillation. None of them worked.
So does anyone know what the minimal concentration for this to work is?
Normally I just distill first and then salt out most of the rest of the water with potassium carbonate, then distill with anhydrous magesium sulfate to achieve ~97% or so.
If I could salt out before the first distillation that would be great.

Zombie - 8-3-2015 at 14:36

Quote: Originally posted by deltaH

Quote: Originally posted by Zombie

You're beating a dead horse brother.

[Edited on 8-3-2015 by deltaH]

You are correct. I miss stated that part Less energy...

Now a Salt Station (thumper) may be worth a try.
Or same thing for 3AZeolite.

It would be easy to spot the point of saturation, and amounts could be modified.

I would think an empty column filled with a desiccant, with a RBF below to collect water May work.
The only issue is the column would have to be heated enough to pass EtOH vapor, and cool enough to allow the desiccant to trap / hold the water.

A simple "dryer" in the vapor path.

You still have the issue of the hygroscopic nature of EtOH collecting moisture from the environment to contend with.

I mean otherwise you just start w/ Everclear, and don't bother drying it.

j_sum1 - 8-3-2015 at 14:45

The link that chem_haruka cited upthread recommended Na2SO4. I didn't read the whole thing but I did skip to the end:

Quote:

from page 38

6. CONCLUSIONS
1. The best of the three new separation designs uses a beer still
to concentrate the feed to 50 wt % solvents, KF salt, and a multipleeffect
evaporator for salt recov-ry.
mental capital costs ($1.72 x 10 vs $1 76 x 10 ) and iiiu h more favorable
incremental operating costs ($2.14 x lot/, vs $4.83 x 10 /y) than the
conventional separation.
This process has comparable incremental operating costs than the conventional separation.
2. Na2SO4, the best salt for this process, effected good phase separation,
while its low water solubility means a low salt addition rate.
Addition of this salt can break both the butanol/water and the propanol/
water azeotropes,
3. Evaporation is better thdn precipitation for salt recovery in
this process, because of the low salt losses and much lower energy requirements
compared with precipitation.
7. RECOMMENDATIONS
1. The use of salt to separate alcoho!/water mixtures is effective,
and further investigation is definitely recommended.
2, For the Clostridia fermentation-product separation specifically,
a parameteric study should be performed to optimize the separation design
presented here.
3. The use of this process in other organic-aqueous separations
should be investigated.

Zombie - 8-3-2015 at 14:52

Quote:

"2. Na2SO4, the best salt for this process, effected good phase separation,
while its low water solubility means a low salt addition rate.
Addition of this salt can break both the butanol/water and the propanol/
water azeotropes,"

careysub - 8-3-2015 at 16:06

Quote: Originally posted by Praxichys

...
Another good point was brought up in the thread regarding the cost of potassium carbonate. It could be made in-situ with super cheap sodium carbonate and KCl....

Does this really work? Sodium and potassium chlorides are about equally soluble in water, and so are sodium and potassium carbonates, so you will not be able to separate the mixture of ions.

The mix might work for salting out the ethanol though, has it been tried? The linked thread raised the possibility but no references or results were posted.

I can get pure sodium bicarbonate (to be dehydrated in the oven) and potassium chloride both for 50 cents a pound, these are by far the cheapest reagents I have access to (other than water).

I had been debating the best way to get cheap reagent ethanol: starting with cheap vodka, or with Kleenstrip denatured alcohol (roughly half and half methanol and ethanol with ~1% MIBK). The Kleenstrip is the cheapest solvent available to me, other than gasoline.

Not sure which would give the cleanest product - if you leave enough tails the MIBK contamination might no be too bad. How bad is the contamination of miscellaneous fermentation products in vodka?

[Edited on 9-3-2015 by careysub]

Zombie - 8-3-2015 at 16:14

That all depends on the brand.

Cheap store bought liquor is mainly what distillers call "Heads". It's the acetone, and Lighter Ketones.

They are lighter, and can be distilled out with careful Vapor Temp. control.

You're best bet for "store bought is EverClear... 190 proof or 95%ABV

http://en.wikipedia.org/wiki/Everclear_%28alcohol%29

BromicAcid - 8-3-2015 at 19:25

Quote: Originally posted by Zombie

To pre separate EtOH from water it is easier to freeze the water , an pour off the EtOH. Due to the volatility you really don't want to begin with more that a 40% ABV in a "still".[Edited on 3-8-2015 by Zombie]

http://www.engineeringtoolbox.com/ethanol-water-d_989.html

Unless you're starting from beer or wine you're not going to get much more concentrated by fractional freezing. By the time you get up to 35% EtOH you're pushing the limits of a household freezer. There are plenty of ways to separate the water from EtOH, each with their own merits depending on what equipment and reagents you have handy. This salting out is somewhat interesting but it is only another tool in the tool box.

careysub - 8-3-2015 at 20:27

Quote: Originally posted by Zombie

That all depends on the brand.

Cheap store bought liquor is mainly what distillers call "Heads". It's the acetone, and Lighter Ketones.

I was hoping that by buying the absolute cheapest vodka (per mL ethanol) at BevMo I would score some that legendary Archer Daniels Midland vodka, the stuff made from industrial 95% ethanol that is bought by the barrel and simply diluted with water.

Vodka drinkers no doubt recoil in horror, but the industrial ethanol is exactly want I want.

Quote:

You're best bet for "store bought is EverClear... 190 proof or 95%ABV

http://en.wikipedia.org/wiki/Everclear_%28alcohol%29

My other option was to go with 75% Everclear, which is the highest proof available in California. Per mL it is about twice as expensive. For that I was going to try salting and then mol sieves.

I wondered about this state ban on 190 proof Everclear (maybe this is just corn-fed industrial ADM hooch, undiluted)? You're okay with 150 proof, but 190 is just too much? Really?

But then I thought that maybe it is a fire safety things. Liquor stores catching on fire is not exactly unknown. Storing a highly flammable solvent in plastic bottles, treating it like a food item in the supply chain, does have its problems.

Zombie - 8-3-2015 at 20:31

Quote: Originally posted by BromicAcid

Quote: Originally posted by Zombie

This is about the most sensible post in the thread. (mine included)

Depending on your starting point, and where you need to end up... There are different routes. :cool:

Zombie - 8-3-2015 at 20:37

Quote: Originally posted by careysub

Quote: Originally posted by Zombie

That all depends on the brand.

Cheap store bought liquor is mainly what distillers call "Heads". It's the acetone, and Lighter Ketones.

Quote:

You're best bet for "store bought is EverClear... 190 proof or 95%ABV

http://en.wikipedia.org/wiki/Everclear_%28alcohol%29

The other thing you have to keep in mind is finished quantity.

Cheep hootch is five bucks per liter.
Everclear Azeotrope is 10 bucks

You're only recovering 35%ish from a liter of cheep stuff.

For experimental purposes I think the idea is fun. For anything really practical... Everclear, and 3A filtration is about the most cost effective way to go. The pellets will out last your lab if you dry, and store them properly.

Just sayin'

[Edited on 3-9-2015 by Zombie]

deltaH - 8-3-2015 at 21:44

Yes, 3A zeolite is a good way to dry azeotropic ethanol. What I'm looking for is the chemical 'treatment' for the step before that, that is to take liquor up to something that you won't have to use a ton of zeolite on. I'm hoping that a salting out combined with a one stage distillation can help here.

Then there's the question of what to do with mash to take it to liquor (beside distillation). As Molecular Manipulations stated, salting out didn't work at such low concentrations.

Zombie has raised the option of freezing as another method of enriching, can this be done with mash... more importantly, does it work well on mash?

What about solvent extraction and what about using cooking oil? At what stage of the enrichment train, if any, could 'solvex' be used? e.g. Could you extract mash with cooking oil, then heat to drive off the vapours and condense strong liquor ready for salting out distillation?

There lots of room for experimentation here

Zombie - 8-3-2015 at 22:06

A single stage distillation can't hit Azeo unless you start w/ azeo.

I'm exaggerating but not by much. Freezing mash is one of the oldest ways of removing water.

The oil thing I have never heard or thought about.

Salting mash sounds impractical.

I do kinda like the idea of an in line drier. It's really a fuel ethanol kinda thing tho.

clearly_not_atara - 8-3-2015 at 23:38

Quote:

Hypothesis:

Saturating drinking gin with anhydrous zinc chloride will increase the alcohol concentration of the vapours enough to make it flammable.

Hypothesis:

ZnCl2 + H2O >> HCl + "ZnClOH"

HCl + EtOH [cat ZnCl2] >> EtCl + H2O

I'm, um, slightly afraid that you might have cancer. Just be glad you didn't distill it! (SN2 is slow at RT)

[Edited on 9-3-2015 by clearly_not_atara]

deltaH - 8-3-2015 at 23:48

Solvent extraction is *just* another kind of phase separation technique, but it works best when there's favourable partitioning between the thing you're trying to extract and the oil to that of it's aqueous phase. With ethanol, this is probably not that favourable, so probably a lousy idea.

My gut feeling still tells me that there might be something in using K3PO4 with intermediate concentration alcohol, I'll be experimenting with that.

Yes, this is more for fuel and solvent ethanol applications than drinking.

********************************************************
Just had a brainwave...

One might combine a solvent extraction with salting out, that is, saturate with Na2SO4, K2CO3 or K3PO4 and also use whatever oil you like for the second phase (gasoline or a paraffin if you're going the fuel route)

From what Praxichys said earlier, one of the problems with salting out is that there's too much water in the organic phase that forms, but if one used an oil there, then you might depress that strongly.

The salt makes ethanol not want to be in the aqueous phase while the oil makes water not want to be in the organic phase.

[Edited on 9-3-2015 by deltaH]

deltaH - 9-3-2015 at 00:06

Quote: Originally posted by clearly_not_atara

Quote:

Hypothesis:

Saturating drinking gin with anhydrous zinc chloride will increase the alcohol concentration of the vapours enough to make it flammable.

I'm not afraid at all of this occurring to any amount worth worrying about, but thanks for the concern.

As for the cancer and since I'm a caucasian in Africa, risk from sun exposure is 10^17 greater (yes I quickly worked that out... not)

careysub - 9-3-2015 at 05:40

Quote: Originally posted by Zombie

I do kinda like the idea of an in line drier. It's really a fuel ethanol kinda thing tho.

Speaking of in-line driers, this discussion is not complete without mentioning to very good performance of warm, dry corn grits in a drying tube.

http://journeytoforever.org/biofuel_library/ethanol_grits.ht...

The grits must be heated to 80 - 110 C so that the process is direct absorption of water vapor by the polysaccharides in the grits, but it can produce 99.9% ethanol. The grits can be take quite a large water load before needing to be replaced/regenerated, and they can be dehydrated and reused repeatedly (probably not worth the trouble for a lab scale process).

[Edited on 9-3-2015 by careysub]

Zombie - 9-3-2015 at 08:55

A few distillers from the forums I frequent have tried corn as a drying agent. All withe similar results.

At or around 98 -99% ABV but all of them had the abv drop to Azeo within a few hours. 6 hours was the longest it held.

Now if you stored it in inert gas or a vacuum sealed container it will hold. This is a whole 'nuther topic tho...

I finally understood the point of the oil... Nice theory, I'd like to see what happens.

deltaH - 9-3-2015 at 09:46

I am not kitted for distillation, but here's the experiment I'm thinking of:

Saturate cheap liquor with potassium carbonate. Extract with three equal volumes of cooking oil (e.g. sunflower) by shaking for a minute, then allowing to separate. Combine oil extract portions into one and distill (the distillation is more of stripping the alcohol because clearly the oil 'ain't going nowhere'). Pour liquor tails with residual salt in a tray and leave out in the sun to dry to recover the salt.

It would be great if someone would be willing to try this and report on what's recovered.

papaya - 9-3-2015 at 11:12

deltaH, I'm not sure vegetable oil is miscible with ethanol, are you sure it is?

deltaH - 9-3-2015 at 11:27

Quote: Originally posted by papaya

deltaH, I'm not sure vegetable oil is miscible with ethanol, are you sure it is?

No I'm not

Perhaps wise to first try dissolving a little methylated spirits into some vegetable oil to check. Thanks for pointing that out.

Well at least we know gasoline would definitely work, but then that's only good for fuel applications. I wouldn't distill the gasoline of course, but maybe dry it if necessary.

[Edited on 9-3-2015 by deltaH]

pneumatician - 22-3-2015 at 09:55

I read in POP-SCI a method to increase the alcohol concentration:

"HIDE GLUE for alcohol dryness" but I never make a try

here a lot of info:

https://archive.org/details/alcoholitsproduc00simmuoft

a tech is put benzene but this is, maybe, the reason of news here and there of explosions in alcohol factories, apart benzene is cancer stimulator so for human use is no-no.

aga - 11-11-2015 at 13:53

deltaH was off on one again while we were discussing this Ethanol separation method on skype.

The notion is that tripotassium phosphate is sufficiently hygroscopic and ethanol-hating that it can rip ALL of the water from the EtOH/Water azeotrope.

Sounded like bollocks to me, however there's One way to find out ...

K₃PO₄ was prepared by reacting 3 mol equivalents of KOH with phosphoric acid.

This is a bitch to get to crystallise and dry, however determined boiling and a couple of hours in an oven @ 130 C worked to yield dry white powder (the hydrated crystal lumps are rock hard, a bit like Al Sulph).

An EtOH refractometer was calibrated with distilled water (DW), then some OTC Vodka was tested to be 37w% EtOH.

The solubility in water of K₃PO₄ calculates to require ~57g to completely saturate the water in 100ml of this Vodka.

60g of dry tripotassium phosphate was loaded into a 250ml conical flask and 100ml of the Vodka was added, with a stirbar running for 4 minutes.

With stirring switched off, this mixture quickly formed 2 layers - within about 15 seconds.

This bi-phasic mixture was put into a 250ml separatory funnel, then quickly dumped back into the conical flask - the undissolved crystals would have blocked the tiny exit hole.

An attempt was made to filter with a standard filter paper, however the rate of flow was stupidly low, and the two phases were seen in the funnel.

Filtering into the sep funnel with a coffee filter was successful.

On deltaH's suggestion, some of the crude mixture was placed in a test tube and two drops of phenolpthalein were added.

... the idea being that the indicator would remain in the upper organic ethanol layer.

Seems he was right about that.

2ml of the Upper layer in the sep funnel were drawn off with a pipette and tested with the refractometer.

The reading said 100w%.

The refractometer was cleaned and re-calibrated with DW (no adjustment was necessary) and the upper layer re-tested.

100 w%

Impossible result, so a volume was drawn off in a 10ml graduated pipette and precisely 3.1ml was mixed with 3.1ml of DW.

This mixture was mixed thoroughly and tested with the refractometer.

50%. Not 49 or 51 or anywhere between. Exactly 50%.

The refractometer was again cleaned and recalibrated.
Again, no adjustment was necessary.

The original Vodka was re-tested, giving a reading of 37%.

Congratulations deltaH, you've done it again !

Crude Ethanol concentration to 100% with no distillation

Got to be a World First.

(please direct all chemistry questions to deltaH as i'm just the spanner-monkey in this experiment)

[Edited on 11-11-2015 by aga]

Outakes

aga - 11-11-2015 at 14:42

Naturally nothing ever goes 100% according to plan.

A 140mm 24/40 ground glass-jointed funnel was used for the filtering.

I assumed the Seal of the ground glass joints was preventing the first filtration attempt, so raised the funnel using a clothes peg.

When operating the sep funnel valve the entire thing fell off, spilling about half of the product all over the desk area.

deltaH - 11-11-2015 at 21:44

WARNING: For those of you who might use it for demonstration, please warn kids that drinking 100% ethanol WILL kill you! Absolute alcohol is solely for technical use, NOT consumption.

STUNNING WORK once again aga! I'm extremely grateful that you volunteered your time on this and thank you!

I must say, I am shocked at the 100% result!

I expected K3PO4 to work better than the other salts used because the anion has a higher charge for a similar size and so a higher charge density than SO4(2-) and CO3(2-), but I would NEVER have expected 100%, something close to azeotrope at best.

Remember that second liquid phase is still in equilibrium with an aqueous phase, so the fact that this is so far skewed as to be 100% is incredible to me as a chemical engineer.

I could have expected this if the solid formed a solid hydrate and so sank most of the water, but not in a million years from an aqueous LIQUID layer in equilibrium. ABSOLUTELY MIND BLOWING!

Seems there was no need for using oil in the second layer to get 100%, still can't believe it WOW WOW WOW!

QUESTIONS:

(1) Is it really necessary to filter? Can't you just let it stand covered for some time and then carefully decant off the alcohol layer?

(2) How well does this work with the equivalent of an alcoholic mash? i.e. ~10% v/v ethanol? On second thought probably not necessary as perhaps it's still a good idea to distill a mash at least once just to clean it up and work with more manageable amounts of K3PO4 and also make it easy to recycle the K3PO4 from a clean solution afterwards.

Well done and thanks once again!

P.S. Glass blowers, we still love you

[Edited on 12-11-2015 by deltaH]

JJay - 11-11-2015 at 21:58

I'm impressed.

aga - 11-11-2015 at 23:40

Quote:

QUESTIONS:

(1) Is it really necessary to filter?

No, not really, just that i was excited and got impatient (hence the spill).

I'll try a smaller amount of beer with a drop of phenolpthalein.

Can anyone out there please repeat the experiment to verify that we're not dreaming ?

[Edited on 12-11-2015 by aga]

aga - 12-11-2015 at 00:30

A few moments ago 43g of dry K₃PO₄ was added to 50ml of 4.5% beer.

The resulting mixture resembles more of a Gel than a liquid.

A few small clear droplets were occasionally observed on the top of the mixture, however the turbulence caused by the CO₂ 'fizz' made observation very difficult.

Likely that this process would not be useful for low EtOH 'mash' concentrations.

JJay - 12-11-2015 at 01:26

I'd be interested in giving it a try, but I don't have any TPP... is it available OTC?

aga - 12-11-2015 at 04:52

Phosphoric acid is available (here) from farming suppliers.

JJay - 12-11-2015 at 05:55

I can get phosphoric acid. I guess I could react that with potassium carbonate.

deltaH - 12-11-2015 at 09:42

Potassium hydroxide and phosphoric acid are readily available across the globe and are low-cost chemicals, so too is MKP fertiliser (monopotassium phosphate). For the latter, you'd need to react it with an additional two equivalents of potassium hydroxide to yield the required K3PO4.

What I'm very happy about is that now the amateur chemist can easily make their own auto fuel. Converting gallon amounts of a simple distillate into absolute is now a trivial matter of 'bucket chemistry'.

Unfortunately, the current low oil price is not exactly helpful in my green crusade, but I'd be very happy if someone ran their scooter or car on alcohol prepared by this method.

Where's zombie when you need him!

[Edited on 12-11-2015 by deltaH]

Fulmen - 12-11-2015 at 10:09

I'm simply stunned by this discovery, well done indeed.

Starcruiser - 12-11-2015 at 10:33

Anhydrous CuSO4 works like a charm for drying azeotropic ethanol to (almost) 100% . Just give it a try.

deltaH - 12-11-2015 at 10:38

This is taking 37 wt.% to 100% in one step, not just breaking the azeotrope. The cost of CuSO4 for such an endeavor would be prohibitive.

[Edited on 12-11-2015 by deltaH]

aga - 12-11-2015 at 11:23

Quote: Originally posted by Starcruiser

Anhydrous CuSO4 works like a charm for drying azeotropic ethanol to (almost) 100% . Just give it a try.

Cool ! Another method !

Please do the experiment and show the results.

It would be great to see a comparison.

I'll replicate your experiment, then you do the same with mine and let's see if wer're onto something here.

[Edited on 12-11-2015 by aga]

Starcruiser - 12-11-2015 at 11:28

I was surely thinking about a small scale experiment (500 ml or less). I`ve done it, it works. Dryied 95% ethanol with CuSO4 anh. then distilled. Go and try it

aga - 12-11-2015 at 11:32

Then Distilled ?

Have you not read the post containing the experiment ?

Please detail exactly what you Did including the drying of the copper sulphate hydrate crystals.

[Edited on 12-11-2015 by aga]

Tsjerk - 12-11-2015 at 11:51

Very nice results! Maybe you could add 10ml of your ethanol to 0.1 gram of blue CuSO4? I know that absolute ethanol will discolour it, while 99-ish ethanol won't

aga - 12-11-2015 at 12:02

How would that work Tsjerk ?

Does the copper sulphate need to be 100% dry as well ?

ISTR that it becomes colourless when anhydrous ( no water=no ions) so it'd not be Blue.

Tsjerk - 12-11-2015 at 12:12

The CuSO4 has to be hydrated (blue). Absolute ethanol will dehydrate it to its white state, but it will only do so if it is really anhydrous by itself, making it a nice control for your sample!

Just put a bit in a tube and add the ethanol to be tested to it. If this works and you want to be really sure you could do the same test with your sample but adding a bit (1%-ish) water to the ethanol before adding it to the sulfate.

Edit: this is knowledge from years and years ago... I could have a look if I can reproduce my memories when I get back to the lab tomorrow. We have absolute ethanol, and probably CuSO4 as well. I'm happy I now feel the freedom to do some of my own experiments again! I even go in in the weekends pretending I'm working....

[Edited on 12-11-2015 by Tsjerk]

Starcruiser - 12-11-2015 at 12:12

Quote: Originally posted by aga

Then Distilled ?

Have you not read the post containing the experiment ?

Please detail exactly what you Did including the drying of the copper sulphate hydrate crystals.

[Edited on 12-11-2015 by aga]

If I`m missing something important, I sincerely apologize (it`s a long topic...).

CuSO4 pentahydrate is readily available and, I think, very cheap everywhere. Drying it is as simple as heating it to 150 C for some time (`till it becomes white-gray). Calculate the amount of H2O needed to get it back to pentahydrate crystallin form and so you will know how much water it can take out from the 95% ethanol. Then, as I said, distill the ethanol, leaving the CuSO4 in the pot.

You can of course dry the pentahydrate again and reuse it countless times.

Perhaps this can be done even with lower conc. ethanol, but I`ve tried it only with 95%.

[Edited on 12-11-2015 by Starcruiser]

Tsjerk - 12-11-2015 at 12:20

What I think Aga is referring to is the fact that his method doesn't need distillation, making it more excessable and less time consuming.

aga - 12-11-2015 at 12:21

Quote: Originally posted by Tsjerk

The CuSO4 has to be hydrated (blue). Absolute ethanol will dehydrate it to its white state

Ah. OK.

Will try that.

Will also try with 98% EtOH and see it the effect is different.

Sure would be great to throw this refractometer away and use perceived crystal colours instead ...

aga - 12-11-2015 at 12:24

Quote: Originally posted by Starcruiser

If I`m missing something important, I sincerely apologize (it`s a long topic...)

Just read at most this page and the one before of the Vastness of the whole 3 pages and maybe you'd have a clue.

This is NOT a 'long' topic.

25 pages would be start to be called 'long'.

The ability to read, re-read and understand is a key part of any learning process.

[Edited on 12-11-2015 by aga]

Starcruiser - 12-11-2015 at 12:38

Quote: Originally posted by aga

Quote: Originally posted by Starcruiser

If I`m missing something important, I sincerely apologize (it`s a long topic...)

Just read at most this page and the one before of the Vastness of the whole 3 pages and maybe you'd have a clue.

This is NOT a 'long' topic.

25 pages would be start to be called 'long'.

The ability ro read, re-read and understand is a key part of any learning process.

[Edited on 12-11-2015 by aga]

I presume It might work even without distillation if this is the challenge. I should just let the anh. CuSO4 in ethanol until it gets fully blue once again. The problem is that I`ve run out for now of 95% ethanol

[Edited on 12-11-2015 by Starcruiser]

deltaH - 12-11-2015 at 12:40

Buy a bottle of vodka, add K3PO4 to it, then add 5% water to the organic layer and you will have lots of 95% ethanol to experiment with again

Starcruiser - 12-11-2015 at 12:43

Quote: Originally posted by deltaH

Buy a bottle of vodka, add K3PO4 to it, then add 5% water to the organic layer and you will have lots of 95% ethanol to experiment with again

That`s wise but I`m afraid I will be tempted to use that vodka for more recreational purposes (though I prefer brandy for that matter

)

aga - 12-11-2015 at 13:01

I think the bit you're missing Starcruiser is that deltaH's process rips all of the water out of Vodka in a way that the Ethanol (EtOH) can be easily separated as 100% (likely near, not exact, +/-1%) from that mixture.

Tripotassium phosphate is also non-toxic, which kinda counts for what people might do with the process.

Copper sulphate IS also insoluble in EtOH, however it's ability to rip the water from the EtOH layer relies almost entirely on the hydration of the copper sulphate, and remains untested.

Do the drying and calculations (everyone here will help if asked) and do the experiment yourself, then post the results.

This is not a Challenge as such, as the only Prize is everybody reading this gains an iota more of Knowledge - we all win no matter how it turns out.

[Edited on 12-11-2015 by aga]

aga - 12-11-2015 at 14:33

Strange.

If I proposed such an idea and were asked to reproduce the results, i'd positively Jump at the opportunity.

Perhaps i'm free-er to actually Do things than most people.

In the face of Silence i guess it's up to me to actually Do something and then post the results for actual discusson.

I Do wish people would Do more chemistry - the results are far more interesting than repeated speculation with no hard data.

softbeard - 12-11-2015 at 15:45

Aga, this really is an amazing result if you can get it confirmed. I don't doubt the results you're getting from your refractometer, but the 100% ethanol value needs to be verified.
I would not expect K₃PO₄ to be able to do this separation, and I don't think the dehydration of ethanol like this to be thermodynamically favourable.

But that's just my own thinking; unexpected results make things interesting! Maybe here's a practical confirmation test: Try to dissolve a few milligrams of potassium permanganate in your ethanol. Anhydrous ethanol should remain colourless. The more colour, the more water. If calibrated, the test could be made quantitative. Using a spectrophotometer would be best, but even a test tube to test tube colour comparison should work well.

deltaH - 12-11-2015 at 21:40

Besides salting out ethanol, I've been trying to think of other applications for K3PO4's salting out capabilities.

I found a paper (attached) that investigated the salting out of ionic liquid solutions from water using K3PO4, which is an impressive feat since it's a fellow ionic compound it's forcing out.

Attachment: Salting out of ionic liquids.pdf (2.8MB)
This file has been downloaded 707 times

Truly K3PO4 must be 'King of Salts'

Perhaps it might be used to regenerate choline chloride based deep eutectic solvents (DES) which got diluted with water?

Or even preparing them when all you have is the readily attainable and low cost 70-75% solutions of choline chloride, i.e. dissolve your ethylene glycol/malic acid or whatever else you wish to use as a hydrogen bond donor or a metal salt hydrate, e.g. ZnCl2, into 70-75% choline chloride solution and then saturate with K3PO4 and decant off the DES.

That would simplify the preparation of DES's for the amateur chemist significantly if it would work.

Who would have thought?

I'm officially renaming absolute alcohol derived from K3PO4 to...

or KA-POW! if you prefer.

[Edited on 13-11-2015 by deltaH]

deltaH - 12-11-2015 at 23:26

For those not yet clear on this, there's a big difference between salting out and desiccation. Using an anhydrous salt that wants to form a hydrate is a type of desiccation. It's a solid-liquid system and absorbs stoichiometric amounts of crystal water.

Salting out is a liquid-liquid or biphasic liquid equilibrium, it is non-stoichiometric and you need much less salt to achieve large changes in concentration.

chemrox - 13-11-2015 at 04:17

anyone know what a "multipleeffect evaporator" is?

Praxichys - 13-11-2015 at 06:46

Unfortunately the cheap phosphate fertilizer I see is ground-up phosphate rock. Buying monobasic potassium phosphate for conversion to tribasic is still like $7 per pound. I don't think this is any more efficient than the potassium carbonate method. You could convert K2CO3 to the phosphate with acid.

Where are you getting your fertilizer?

aga - 13-11-2015 at 09:14

Unfortunately i dumped the dried ethanol from the previous experiment into my ethanol reserve, so to test some ideas i made a small qty more with the remaining K₃PO₄.

The same bi-phasic phenomenon was observed with the ethanol quickly moving to the Upper layer.

Here is the Upper layer with some KMnO₄ as suggested by softbeard with the right hand tube containing distilled water + postassium permanganate for comparison.

Whilst cleaning up (~1 hour) it was noted that the KMnO₄ had fallen out of suspension and was only seen on the bottom of the test tube.

As suggested by Tsjerk a single CuSO₄.5H₂O crystal was covered with the dried ethanol.

No immediate colour change was noted in either the liquid or the crystal.

After 15 minutes a colour change had occurred in the crystal, but not the liquid.

Next on to the suggestion made by Starcruiser, which was rather more arduous.

100g of copper sulphate crystals was dried in an oven at ~150 C for two hours, with periodic removal, crushing and spreading of the crystals, yeilding a not-quite-white powder (still had a slight green tinge).

The resulting drier copper sulphate weighed 68.86g, representing approximately 86% of the crystal water removed.

10ml of the the 37% Vodka was diluted with 10g of solidified tap water. To this was added around 20ml of a citrate extraction.
This was mixed, then added to the amateur chemist

Next 80ml of the Vodka was measured with a cylinder and poured into a 100ml beaker.

Theoretically the water in 80ml should be absorbed by 28g of anhydrous copper sulphate, so 28g was added as a starting point, and mixed thoroughly.

The powder immediately became Blue again and slight heating was noted.

After 10 minutes nothing remarkable was seen, so more of the dried powder was added and mixed.

After adding a further 10g, there was no room to add more.

Some of this blue liquid was tested in the refractometer, and measured at 46w% EtOH.

This is a light-based instrument, so it is highly likely that the blue colouration affected this result.

To test the Colour effect, a small quantity of the dried powder was dissolved in distilled water to match the colour of the liquid in the beaker as close as practicable by eye.

This measured 7w% despite there being no EtOH in there.

Made reckless by the Vodka & Orange from earlier, i set light to the liquid in the beaker.

It burned merrily with a colourless flame, so it is likely that some degree of water sequestration had happened.

Overall the Anhydrous Copper Sulphate method left the Ethanol layer with a significant quantity of water, as the Blue colour testifies (no water, no ions, no Blue).

By comparison, the dried copper sulphate would not dissolve at all in the Ethanol produced by the tripotassium phosphate method.

Please deride, slate and comment.

deltaH - 13-11-2015 at 09:31

Very thorough work again aga, well done!

deltaH - 13-11-2015 at 09:39

Quote: Originally posted by chemrox

anyone know what a "multipleeffect evaporator" is?

Yes, it's standard industrial equipment run as a series under reduced pressure that gets incrementally lower and lower in the evaporators. In a nutshell, it's a way of evaporating a liquid to concentrate a solution in a more energy efficient way than doing it in just one heated step. You use the vapour from a previous step at a higher pressure to boil the liquid in the next step and so convert some of the heat of condensing the vapour into boiling the liquid. Overall you still need net heat, but by recovering the heat of condensing the vapour, you save on a lot of energy.

Read more here: https://en.wikipedia.org/wiki/Multiple-effect_evaporator

deltaH - 13-11-2015 at 09:47

Quote: Originally posted by Praxichys

I don't think this is any more efficient than the potassium carbonate method.

It would be interesting to see what percentage alcohol results when using potassium carbonate.

We already know that it does form a second liquid layer from the opening post's reference, just not what concentration results.

If it too makes an absolute alcohol, then sure, it would be preferable.

[Edited on 13-11-2015 by deltaH]

aga - 13-11-2015 at 11:30

Ah crap.

OK. So do the same with potassium carbonate.

Think i'll just buy that reagent this time.

deltaH - 13-11-2015 at 11:59

It's solubility is 112g/100ml water, so you could use a little more to saturate the solution than what you did with K3PO4.

[Edited on 13-11-2015 by deltaH]

aga - 13-11-2015 at 12:55

One would do a calculation or two in any event.

Thanks for looking up the solubility.

aga - 13-11-2015 at 13:23

Before i get going with any more experiments, are there any OTHER substances to test in this application ?

May as well get the whole series sorted in one go.

MolecularWorld - 13-11-2015 at 17:21

You could try trisodium phosphate.

I wanted to try this tripotassium phosphate process, but I don't have alkaline potassium compounds or phosphoric acid.
What I do have is trisodium phosphate, available as a cleaning product* at many hardware stores.
I also don't have vodka or similar ethanol solution, but I did have some old cheap whiskey (40% ABV), and 40% methanol solution (winter windshield washer fluid).

The trisodium phosphate appears to work for separating ethanol from solution. I lack the means to determine the concentration of the top layer, but it burned readily with a blue flame, leaving a small amount of carbonaceous residue that smelled of burnt sugar. The volume was close to what would be expected in a complete separation.
It also appears that trisodium phosphate is unsuitable for separation of methanol from solution. There is no obvious separation. The white layer on the bottom is undissolved trisodium phosphate. Liquid extracted from the top only burned once heated to boiling, and evaporated to leave sodium-containing residue.

*This is being phased out in favor of misleadingly labeled TSP_substitute. Be sure your TSP is real, the substitutes are useless.

[Edited on 14-11-2015 by MolecularWorld]

UC235 - 13-11-2015 at 20:33

Quote: Originally posted by MolecularWorld

You could try trisodium phosphate.

I wanted to try this tripotassium phosphate process, but I don't have alkaline potassium compounds or phosphoric acid.
What I do have is trisodium phosphate, available as a cleaning product* at many hardware stores.
I also don't have vodka or similar ethanol solution, but I did have some old cheap whiskey (40% ABV), and 40% methanol solution (winter windshield washer fluid).

The trisodium phosphate appears to work for separating ethanol from solution. I lack the means to determine the concentration of the top layer, but it burned readily with a blue flame, leaving a small amount of carbonaceous residue that smelled of burnt sugar. The volume was close to what would be expected in a complete separation.
It also appears that trisodium phosphate is unsuitable for separation of methanol from solution. There is no obvious separation. The white layer on the bottom is undissolved trisodium phosphate. Liquid extracted from the top only burned once heated to boiling, and evaporated to leave sodium-containing residue.

*This is being phased out in favor of misleadingly labeled TSP_substitute. Be sure your TSP is real, the substitutes are useless.

[Edited on 14-11-2015 by MolecularWorld]

Check your MSDS too. "TSP" as sold is frequently just a mix of TSP and sodium sesquicarbonate with traces of sodium silicate. I think they neutralize crude phosphoric acid produced from sulfuric acid and calcium phosphate with trona in excess, dry the result, and throw it in a box. It is far from pure, but not a bad starting point.

In fact, here is the MSDS from the page you linked. http://www.homedepot.com/catalog/pdfImages/5b/5b08c5e5-d610-...

[Edited on 14-11-2015 by UC235]

MolecularWorld - 13-11-2015 at 21:03

The TSP I used above is the same brand as at the link.
I never said it was pure, and my experiment shows it doesn't have to be.
I didn't even dry it before use.
I read the MSDS (how'd you think I knew it was real?), and as you said, there is some "sodium sesquicarbonate", but at a low enough concentration that it could still be called technical-grade trisodium phosphate.

The substitutes are completely different. They contain no phosphate (the phase-out is supposedly to reduce phosphates in runoff), and may be any number of other compounds (I almost bought one that turned out to be mostly "sodium metasilicate").

[Edited on 14-11-2015 by MolecularWorld]

deltaH - 13-11-2015 at 21:25

Thank you MolecularWorld for your contribution and work!

I have one more potential trick up my sleeve in regards to this... salting out dilute ammonia solutions.

You almost certainly wouldn't make a second liquid layer, but instead, lots of ammonia vapour which you could lead via a pipe and into ice to make concentrated and pure ammonia solution. The ice is very important to help with absorption since this is very exothermic for ammonia.

OK, I know it's not that impressive, but still, if you need 25% ammonia solution and all you can get is store bought scrubs ammonia solutions contaminated with detergent...

[Edited on 14-11-2015 by deltaH]

deltaH - 13-11-2015 at 21:37

aga, there are a great many other salts that could be tested. In my mind I have the following hypothesis in terms of efficacy:

K3PO4 > K2CO2~Na3PO4~K2SO4~(NH4)2SO4 > Na2CO3~Na2SO4~KCl~NH4Cl > NaCl

... but for the sake of brevity, you might only want to try a select few.

[Edited on 14-11-2015 by deltaH]

MolecularWorld - 13-11-2015 at 21:40

I just saturated some crappy dollar store ~2% ammonia solution with crappy hardware store trisodium phosphate.
The results were inconclusive.

Very slight effervescence was seen.
The ammonia odor was, if anything, reduced, suggesting the ammonia is reacting with either the phosphate or carbonate.

Someone with pure phosphates and somewhat stronger (~5%) ammonia solution should attempt this to allow for better observations.

deltaH - 13-11-2015 at 21:45

Thanks MolecularWorld. I don't think you can have any carbonate in there though, that would sink ammonia to form carbamate anions (NH2COO-) in a fast equilibrium.

You really need relatively pure and anhydrous salts, carbonate free.

deltaH - 13-11-2015 at 22:00

Just read about the terrible attacks in Paris, my condolences to all our French SM members, how terrible

deltaH - 14-11-2015 at 01:56

Chatting to aga, I have had a speculative brain wave... adding K3PO4 to concentrated ammonia solution (34%) might result in extreme cooling.

Why?

The idea is that the K3PO4 dissolves in the conc. ammonia solution, interacts strongly with the water there making the ammonia not want to hang around anymore. The ammonia then boils off, but this is highly endothermic (ammonia has a high heat of vaporization) and so the remaining solution gets colder and colder. Kind of like boiling liquefied gases rapidly out of a gas cylinder causing the cylinder to get colder and colder. If we're REALLY lucky and this goes to extremes, something amazing just might happen, but I don't want to say what just yet because the chances of it happening are pretty slim, anyway, some of you might guess...

Unfortunately one would need a really good fume cupboard for this because of the large amounts of ammonia fumes generated.

[Edited on 14-11-2015 by deltaH]

Tsjerk - 14-11-2015 at 03:11

Unfortunately I only remembered about this experiment (absolute ethanol/ CuSO4 pentahydrate) the moment I went home yesterday. I did start the experiment but also I didn't see an immediate change of color. The tube is sealed air tide, and I will have a look at it Monday morning.

The ethanol I used comes in as 99,9%+ P.A. absolute ethanol, so not completely anhydrous. Also people don't care about moisture from the air coming in because we don't need it to be anhydrous. But I think it should still be close to anhydrous as the bottle was pretty new. I will let you know, I also took a picture before and after adding the ethanol.

softbeard - 14-11-2015 at 13:21

Quote: Originally posted by aga

Unfortunately i dumped the dried ethanol from the previous experiment into my ethanol reserve, so to test some ideas i made a small qty more with the remaining K₃PO₄.

The same bi-phasic phenomenon was observed with the ethanol quickly moving to the Upper layer.

Here is the Upper layer with some KMnO₄ as suggested by softbeard with the right hand tube containing distilled water + postassium permanganate for comparison.

Hey Aga, that potassium permanganate dissolution test looks really convincing! If I remember right, KMnO₄ in 95% ethanol gives a barely perceptible pink colour. Looks like in your test you just got a trace of Mn₃O₄ suspension, no KMnO₄ dissolving at all.
Maybe do a comparison test with the potassium permanganate in one test tube with 95% ethanol and the other tube with potassium permanganate and the K₃PO₄-dried ethanol.

BTW: I think it's great you doing these experiments and showing the pictures! I'll be trying some ethanol drying with the K₃PO₄ when I get some time. Phosphoric acid, I got. I think it's time to look for some K₂CO₃ on eBay.

aga - 14-11-2015 at 15:30

All i got is a fully equipped Shed and the Will to spend about 4 hours Doing Stuff.

Most people have a phone that does photos these days.

Unfortunately i stuck a contaminated 10ml pipette in the Ethanol reserve (250ml bottle) and now it's all milky.

Sodium thiopsulphate maybe, maybe some silver chloride, not sure.
Needs distilling again certainly.

Anyone can Do experiments, and more people Should.

Get some potassium carbonate and Test this thing !

[Edited on 14-11-2015 by aga]

deltaH - 14-11-2015 at 21:23

softbeard, be aware that K3PO4 is a foaming agent and also that these solutions become viscous, so not so sure K2CO3 would be the best feedstock to use, I'd use KOH as aga did, taking great care, of course.

I see on alibaba.com that there are many sellers selling anhydrous food grade K3PO4, abbreviated as ATKP. It's a mere $1-2/kg in bulk quantities, if only it were that easy :mad:

aga - 21-11-2015 at 09:16

Quote: Originally posted by deltaH

Chatting to aga, I have had a speculative brain wave... adding K3PO4 to concentrated ammonia solution (34%) might result in extreme cooling.

It seems to result in extreme ammonia production at least !

A cork was drilled to accept some glass piping and a glass thermometer.

A plastic tube was attached to the glass pipe and connected to an inverted funnel to serve as anti-suckback in a 250ml beaker containing iced water.

10g of dried K₃PO₄ was loaded into a 25ml RBF.

10ml of 33w% aqueous NH₃ was injected into the RBF via a hypodermic syringe via the take-off pipe in two portions (the biggest hypodermic on hand was 5ml).

A reaction was noticed immediately on adding the first few drops of ammonia.

Unfortunately the gas evolution from the flask prevented further ammonia being added (the liquid kept bubbling up the pipe, and not into the RBF) so the cork was released, and the ammonia injected.

On re-seating the cork, strong bubbling of ammonia was seen in the scrubbing beaker.

The RBF became cold to the touch, although not freezing.

The scrubber liquid was tested with UI paper and is strongly basic.

[Edited on 21-11-2015 by aga]

aga - 21-11-2015 at 12:24

Quote: Originally posted by MolecularWorld

You could try trisodium phosphate.

Hmm.

Comparing the effectiveness of tri-K and Na Phosphate sounds like a great idea.

With any luck, may get that done tomorrow, depending on how well the product of H₃PO₄ + NaOH dries.

For the last K attempt i found a stainless steel dog bowl to be most effective.

Comparison of Tri Potassium Phosphate and Potassium Carbonate

aga - 16-12-2015 at 11:17

This experiment was intended to quantitatively compare the salting-out efficiency of the two chemicals in an ethanol/water separation.

15ml Ethanol was diluted with 35ml of distilled water to form the 'feedstock'.

This concentration was tested with a refractomer to be 31 w% <sup>(1)</sup>.

Two test tubes were filled with ~6ml of the feedstock.

To one was added 1.20g dry (homemade) K₃PO₄

The other had 5.47g of K₂CO₃ added.

These amounts were determined by first weighing the reagents, then adding spatulas of the reagent to the test tube, shaking, and repeating until no further reagent would dissolve.

The reagents were then re-weighed, the difference in weight being the quantity used.

(this simple technique is called accurately adding stuff to other stuff to see what happens)

The test tubes were allowed to stand.

Phase separation occurred in both test tubes, with the K₃PO₄ sample achieving stable phases after 30 seconds.

The K₂CO₃ sample did not heat very much at all, and took ~8 minutes to achieve stable phases.

The K₃PO₄ sample created 3 phases, presumably due to either insufficient reagent, or (highly likely) a lot of impurities in this home-made batch.

Presuming the upper layers to be in the 90+ w% range, direct reading would not be useful with the refractometer as the scale is logarithmic, with the lower part of the scale being more accurate.

So 1ml of each liquid phase was drawn off in a pipette and mixed with 1ml DIW.

This 50% diluted solution was then tested, with the result doubled to arrive back at the actual w%.

The Results were :-

Upper K₃PO₄ layer : 56w%
Lower K₃PO₄ layer : 78w%
Upper K₂CO₃ layer : 100w%

Conclusions:-
Pure potassium carbonate works very well in the salting out of ethanol, although is much slower than tripotassium phosphate.

The experiment needs to be repeated with equally pure reagents to determine the exact times to achieve the effects, the heating generated and the actual quantities of ethanol recovered.

<sup>(1)</sup>
Yes, 31w% = 103w% in the original ethanol. This is not a 100% accurate analytical lab ...

Please deride, slate and comment.

deltaH - 16-12-2015 at 12:08

Interesting that the K2CO3 also gave 100%! I'm amazed this isn't more common knowledge.

Looks like you have some metallic contamination in this K3PO4 batch this time (not that it should matter much) and I suspect you didn't use enough of it, hence the lower alcohol purity.

Thanks again for your contribution and nice work!

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