Sciencemadness Discussion Board - Success Separating Methylated Spirits?

Success Separating Methylated Spirits?

Hexavalent - 14-10-2012 at 10:01

In the UK our methylated spirits/denatured alcohol is purple, so I set up a fractional distillation apparatus and begin heating my mixture. As expected, it came over colourless but what surprised me the most was the boiling points - it hung between 64 and 67 degrees Celsius for a while before climbing to 75.5 degrees Celsius where it remained until I shut down the hotplate (with about 5% liquid remaining in the flask).

The 64-67 is indicative of the methanol content and the latter of ethanol....it seems that I may have done some sort of a separation, despite the commonly-held belief that methylated spirits is supposed to be extremely hard to separate. The MSDS for this product is rather vague, listing just it's components - MeOH, EtOH, pyridine and the dye. The last two are likely to be in small amounts, according to other products' trends.

I did no heating control as such, merely switching on my kitchen hotplate to full capacity before lowering it to 75% of its capacity. I was using an air-bath with aluminium foil insulation, and a stainless steel gauze/mesh-type packing for my fractionating column (15 cm effective length).

My thermometer is a mercury one, -20*C - 150*C and I was using an all-glass 19/26 distillation rig.

Any ideas about what happened here? Did I achieve separation, even to some degree?

[Edited on 14-10-2012 by Hexavalent]

hissingnoise - 14-10-2012 at 13:06

Quote:

Any ideas about what happened here? Did I achieve separation, even to some degree?

To some degree, yes, but the distillate may still contain pyridine along with residual CH<sub>3</sub>OH!
To obtain the correct B.P., the mercury bulb <i>must</i>, as you know, be aligned with the adapter side-arm . . .

Hexavalent - 14-10-2012 at 13:23

Yeah, I get that as well as the fact that the final product is still likely to be impure - my main objective was to remove the colourant, and I was surprised by how the temperatue hung around the methanol range for a while before going up to the standard ethanol/water azeotrope range.

P.S. I am absolutely certain that my thermometer was positioned correctly, I spent 5 minutes adjusting it until I got it just right.

12AX7 - 14-10-2012 at 14:13

The pyridine could be ionized with some HCl, which should keep it in the pot, no?

Tim

hissingnoise - 14-10-2012 at 14:21

Jeeez young Tim, long time no hear! Welcome back!

blogfast25 - 15-10-2012 at 09:35

Personally I think your lower initial head temperature was just a warming up effect. Did you run the column at total reflux for at bit, to stabilise temperatures? Had you done that I think you'd have gotten the ethanol-water azeotrope immediately after switching to partial reflux.

Interesting idea to neutralise the pyridine with HCl but any excess is likely to result in an acidulated end-product. Also: not great for your steel packing!

Scr0t - 15-10-2012 at 11:30

I've distilled similar compositions many times.

The dye (methyl violet) and pyridine aren't a problem when distilling however...

Whether it is listed or not it likely contains naphtha or some petroleum distillate and will give a turbid suspension on dilution with H2O even after several fractional distillations.

You will easily be able to achieve some degree of separation of the MeOH and EtOH through this means but if you use it make Et2O from it then it will likely still have an obvious contamination with MeOEt in spite of your best efforts to remove the MeOH.

Hexavalent - 15-10-2012 at 12:11

I've dried the product with MgSO₄ to remove excess water, and when diluted with additional water I get no turbidity...

AJKOER - 15-10-2012 at 12:42

Here is a crazy and expensive path for those who would like to experiment. Add Oxalic acid (toxic see http://www.sciencelab.com/msds.php?msdsId=9926346 ) which is said to readily form Oxalate esthers. Now, as the dimethyl oxalate is a solid below 50 C with a bp of 163.5 (see http://www.sigmaaldrich.com/catalog/product/ALDRICH/135623?l... ) and the diethyl oxalate is a liquid with mp of -41 C and a bp of 185.7 C (see http://www.sciencelab.com/msds.php?msdsId=9923751 ), filter to separate out the Methyl ester and distill with NaOH to recover the Ethanol which has a bp of 78.5 C (see MSDS at http://www.sciencelab.com/msds.php?msdsId=9923955 ). Normally, the ester reaction is reversible and an acid catalyst is employed (see page 45 at https://docs.google.com/viewer?a=v&q=cache:FBgeovLIqXMJ:... ), but in this case, as the dimethyl oxalate separates out, the reaction with Methyl alcohol may proceed without the acid catalyst. If the diethyl oxalate remains in an equilibrium reaction, then recovery of the Ethanol may proceed more easily. Also, Pyridine may also form a separable Oxalate salt.

On this hypothetical expensive process (yes, it has to be cheaper to buy the Ethanol), no guarantee on taste or toxicity (animal experiment perhaps, or your in-laws).

[Edited on 15-10-2012 by AJKOER]

DJF90 - 15-10-2012 at 15:49

@AJKOER: What makes you think that dimethyl oxalate is not soluble in the diethyl ester?

As for the distillation... I've done it before myself. Some 50% H2SO4 was added to keep any pyridine in the stillpot, and heating was commenced. I noticed no methanol come over, despite careful fractionation. The ethanol came over as expected, however the it still had a perculiar smell. I noticed some oily droplets in the stillpot, and concluded it was the naptha causing problems.

I gave up at this point. I used the ethanol in an ether synthesis. The naptha proved easily separable from the diethyl ether, by simple distillation.

AJKOER - 15-10-2012 at 18:18

DJF90:

Per Watts' Dictionary of chemistry, revised and entirely rewritten by Henry Watts, Matthew Moncrieff Pattison Muir, Henry Forster Morley (see page 649, link: http://books.google.com/books?pg=PA650&lpg=PA655&id=... ) one must heat the anhydrous Oxalic acid with the primary alcohol forming a mixture of both the mono- and di-alkyl oxalate. However, upon distilling the mono-alkyl oxalate, there is a decomposition into CO2 and the alkyl formates. So rapid boiling could destroy some of the alkyl oxalate, but perhaps prolonged mild heating may still create the insoluble dimethyl oxalate especially avoiding an excess of CH3OH in which it is soluble (see this link which also notes that (CH3)2C2O4 decomposes in water: http://www.chemicalland21.com/industrialchem/solalc/DIMETHYL... ) by employing an excess of Oxalic acid and adding the spirits slowly with mixing to the dry H2C2O4 over a hot water bath. But still, I agree with you, one should most likely give-up on this procedure especially in the presence of naptha.

[EDIT] Good news, per this link (http://www.mark-ju.net/juliette/meths.htm ), my impression is that the taste poisoning additives and alcohol strength vary considerably by country.

[Edited on 16-10-2012 by AJKOER]

paulr1234 - 15-10-2012 at 21:18

By any chance, does anyone here know a USA based source for purple methylated spirits as described above? It is sold in the UK, Australia and NZ. I'm finding it hard to buy it here. There was this one online source a while back that specialized in Caribbean goods but they seem to have gone of out business (yes, the color and exact smell matters for reasons I won't bore you with).

blogfast25 - 16-10-2012 at 07:22

AJ:

It'd be cheaper (and easier and safer) to ferment sugar to 'wine', filter or decant off the yeast and distill the 'wine' to the alcohol/water azeotrope. Decent wine making yeasts deliver 13 % ABV in about 7 -10 days.

[Edited on 16-10-2012 by blogfast25]

ElectroWin - 16-10-2012 at 09:02

Quote: Originally posted by Hexavalent

As expected, it came over colourless but what surprised me the most was the boiling points - it hung between 64 and 67 degrees Celsius for a while before climbing to 75.5 degrees Celsius where it remained until I shut down the hotplate (with about 5% liquid remaining in the flask).

The 64-67 is indicative of the methanol content and the latter of ethanol....it seems that I may have done some sort of a separation, despite the commonly-held belief that methylated spirits is supposed to be extremely hard to separate. The MSDS for this product is rather vague, listing just it's components - MeOH, EtOH, pyridine and the dye. The last two are likely to be in small amounts, according to other products' trends.

[...]

Any ideas about what happened here? Did I achieve separation, even to some degree?

Yes, it is just what you should expect to happen.

methanol makes an azeotrope with ethanol, which comes off first. then later you get the remaining ethanol; but the loss of yield is considerable.

when commercial distilleries do this, which they need to do because fermentation of pectin produces methanol, they catalyze the oxidation of the methanol to formaldehyde, which makes no azeotrope with ethanol. it also comes off first, but ethanol yield is much better.

UnintentionalChaos - 16-10-2012 at 10:42

Quote: Originally posted by ElectroWin

[citation needed]

There is no methanol-ethanol azeotrope. It's a zeotropic mix. A fractionating column with sufficient theoretical plates can completely seperate the mixture, but cleanly separating liquids with 13C bp differences is no easy task.

That stainless steel packing that you used is quite efficient, it seems. I still would not drink the resulting ethanol, but you could collect a fraction and refractionate it, discarding forerun and the beginning of the main fraction and the ethanol would probably be good enough for most lab purposes.

[Edited on 10-16-12 by UnintentionalChaos]

Hexavalent - 16-10-2012 at 10:44

Can anyone confirm the MeOH/EtOH azeotrope?

tetrahedron - 16-10-2012 at 11:33

what chaos said

Hexavalent - 16-10-2012 at 13:25

Thanks for your help, everyone ~ UC, as I always go by I never consume anything made or that has come from the lab, and my ethanol is indeed just for general lab purposes. In fact, it may soon be used to attempt your ethyl acetoacetate synthesis if I have the time

Another question - can it be predicted whether a ternary azeotrope will form between three compounds, or is it just a matter of reading the literature, looking at tables etc?

Thanks in advance.

tetrahedron - 16-10-2012 at 13:47

ethanol/methanol/water also has no azeotrope, if that's the one you're thinking of ;p

AJKOER - 17-10-2012 at 07:03

Quote: Originally posted by blogfast25

AJ:

It'd be cheaper (and easier and safer) to ferment sugar to 'wine', filter or decant off the yeast and distill the 'wine' to the alcohol/water azeotrope. Decent wine making yeasts deliver 13 % ABV in about 7 -10 days.

Blogfast25:

I agree with your comments, however, a word of caution to the novice. To quote from the "The Home Distillation of Alcohol" (see http://homedistiller.org/intro/methanol/methanol ) "Methanol is formed when fermenting beverages high in pectins - eg grapes and berries. Starting with a grain or sugar based wort, in a clean fermentor with a yeast culture from a well aereated source will result in small/none formed."

hissingnoise - 17-10-2012 at 07:17

And there's always the metathesis of ethyl acetate by a base . . .

tetrahedron - 17-10-2012 at 07:57

the methanol doesn't come from fermentation but from hydrolysis of pectin

blogfast25 - 17-10-2012 at 09:44

AJ:

The amount of MeOH must be negligible: like many I've been making and drinking home made wine for years. I know that site and it's pretty crappy, BTW...

AJKOER - 17-10-2012 at 12:57

More per the same site (http://homedistiller.org/intro/methanol/methanol ):

"Mike explains about the pectin ..
The methanol comes from the pectin, which mainly composed of methyl esters of galactose. When pectin breaks down, by enzymes introduced by microorganisms, or deliberately introduced, the methyl esters combine with water to produce methanol, so the aim should be to leave the pectin well alone if you can.

I think Jack would agree that what he means is that fermenting at a high temperature, or adding pectin enzyme, or trying for an abv higher than 12% all increase the risk of methanol being produced, so his advice about low temperature fermentation, adding no exra enzymes, and a target lower than 12% abv is all good stuff."

Here is another discussion that attempts to quantify the amount of methanol exposure, which appears to be minute, in agreement with Blogfast25, but conditions (type of fruit, how ripe is the fruit processed, high temperatures, excess pectin enzyme and elevated target alcohol concentration) along with the amount consumed may alter this assessment. To quote (link: http://yarchive.net/med/methanol_poisoning.html ):

"As guidance, pectic substances (the major source of methanol in plants)
amounts to 0.5 to 1.0% of the fresh weight of apples. So a 200 gram apple
has 1 to 2 grams of pectin. Pectin starts out about 80-90% methylated, then
during fruit ripening becomes about 40% methylated, because internal plant
enzymes do the cleavage. I'm going to assume that the methyl groups are
about 5% of the total pectic substance (a data-free analysis), so that makes
the 2 g of pectin in an apple have about 50-100 mg of potential methanol.
Assume that only half of these methyl groups are cleaved, so now we're down
to 25-50 mg of methanol in one single apple.

Pectin, and it's de-esterifying enzymes, are found in nearly every fruit,
from strawberries through to bananas. Apple juice would be especially high
in methanol, since manufacturers treat it with commercial pectin esterase to
get rid of cloudiness, and in this case there isn't anywhere for the
methanol to go except to stay in solution.

Although pectin isn't degraded by mammalian enzymes in the human gut, this
doesn't matter because plants themselves contain the enzymes. Moreover,
bacteria in our large intestines will be able to cleave pectin in most
cases, so any residual methylated pectin will form methanol in the bowel.

Really, the quantitative difference between the amount of methanol that
causes blindness in chronic alcoholics who drink it as a cheap drunk, and
the amounts of methanol in either apple juice or aspartame, is so great that
it's really a scam to pretend that there is no difference."

[Edited on 17-10-2012 by AJKOER]

ElectroWin - 17-10-2012 at 13:46

Quote: Originally posted by AJKOER

Really, the quantitative difference between the amount of methanol that
causes blindness in chronic alcoholics who drink it as a cheap drunk, and
the amounts of methanol in either apple juice or aspartame, is so great that
it's really a scam to pretend that there is no difference."

i think the point is, that if you make 300 gallons of moonshine and forget that the first few jugs coming off the still should be discarded, then someone will get hurt.

tetrahedron - 17-10-2012 at 13:59

what this boils down (or..distils) to is that the methanol found in fermented fruit is of no concern; it's only an issue when deliberately added to ethanol (either in a legitimate way as a denaturant, or as an illegal additive, much like melamine in chinese milk). of course, here we wanna do chemistry with our ethanol, not drink it or burn it, so purity matters. it's generally accepted that methanol is impossible to completely remove from ethanol in an amateur setting. look for 'spirits' that don't use methanol as denaturant.

AJKOER - 4-11-2012 at 22:52

Here is a new idea for the general separation of C2H5OH and CH3OH that may be of academic interest/discussion. I would call it the Chlorate method which takes advantage of the reducing properties of Methanol to remove/dilute its presence. More specifically, I will first discuss the reaction with Sulfuric acid that does not work for this separation synthesis. The reaction is reported as:

3 NaClO3 + 2 H2SO4 + .85 CH3OH --> 3 ClO2 + Na3H(SO4)2 + H2O + .o5 CH3OH + .6 CHOOH + .2 CO2

Source: See page 71 at http://books.google.com/books?id=924wS3TfCdIC&pg=PA71&am...

The reaction product is partially confimed at http://www.ahlundberginc.com/chlorine_dioxide_processes.htm . To quote: "Methanol is the reducing agent in the R8 process. The by-product is sodium sesquisulfate [Na3H(S04)2], an acid salt."

Problem with this synthesis are many including the creation of Formic acid, although this can be removed by employing a slight excess of H2SO4 (with the accompanying liberation of carbon monoxide). The generated ClO2, also a problem (explosive gas), can, however, be readily dissolved in water:

ClO2 + H2O <--> HClO2 + HClO3

as part of a chlorate regeneration strategy. However, the reason that the synthesis does not work for the separation of Methanol and Ethanol with Chlorate/Sulfuric acid is that H2SO4 readily (actually exothermic) reacts with the alcohols presence. One possible solution, use HCl which otherwise requires a ZnCl2 catalyst to react with the alcohols. Upon avoiding an excess of HCl, my speculated reaction (unbalanced) is similar to the sulfuric reaction with the additional products of Cl2 and perhaps CO:

NaClO3 + HCl + CH3OH --> ClO2 + Cl2 + NaCl + H2O + CH3OH + CHOOH + CO2+ CO

where I continue to assume that the CH3OH acts, given the limited HCl employed, as a reducing agent as well. In addition, it is assumed that a reaction at room temperature, does not result in a significant loss of the ethanol.

[Edited on 5-11-2012 by AJKOER]

weiming1998 - 5-11-2012 at 00:29

Quote: Originally posted by AJKOER

Here is a new idea for the general separation of C2H5OH and CH3OH that may be of academic interest/discussion. I would call it the Chlorate method which takes advantage of the reducing properties of Methanol to remove/dilute its presence. More specifically, I will first discuss the reaction with Sulfuric acid that does not work for this separation synthesis. The reaction is reported as:

3 NaClO3 + 2 H2SO4 + .85 CH3OH --> 3 ClO2 + Na3H(SO4)2 + H2O + .o5 CH3OH + .6 CHOOH + .2 CO2

Source: See page 71 at http://books.google.com/books?id=924wS3TfCdIC&pg=PA71&am...

The reaction product is partially confimed at http://www.ahlundberginc.com/chlorine_dioxide_processes.htm . To quote: "Methanol is the reducing agent in the R8 process. The by-product is sodium sesquisulfate [Na3H(S04)2], an acid salt."

Problem with this synthesis are many including the creation of Formic acid, although this can be removed by employing a slight excess of H2SO4 (with the accompanying liberation of carbon monoxide). The generated ClO2, also a problem (explosive gas), can, however, be readily dissolved in water:

ClO2 + H2O <--> HClO2 + HClO3

as part of a chlorate regeneration strategy. However, the reason that the synthesis does not work for the separation of Methanol and Ethanol with Chlorate/Sulfuric acid is that H2SO4 readily (actually exothermic) reacts with the alcohols presence. One possible solution, use HCl which otherwise requires a ZnCl2 catalyst to react with the alcohols. Upon avoiding an excess of HCl, my speculated reaction (unbalanced) is similar to the sulfuric reaction with the additional products of Cl2 and perhaps CO:

NaClO3 + HCl + CH3OH --> ClO2 + Cl2 + NaCl + H2O + CH3OH + CHOOH + CO2+ CO

where I continue to assume that the CH3OH acts, given the limited HCl employed, as a reducing agent as well. In addition, it is assumed that a reaction at room temperature, does not result in a significant loss of the ethanol.

[Edited on 5-11-2012 by AJKOER]

Chloric acid is extremely oxidizing, and would oxidise ethanol to acetic acid. I don't think methanol is much more reducing than ethanol, so not all methanol is going to be destroyed, but a lot of ethanol is going to be destroyed. Chlorine dioxide, a product of the reaction, also oxidizes organics (a yellow solution of it readily discolours in the presence of ethanol or glycerol, especially when heated).

A second problem is the violence of such a reaction. I have had mini-explosions occur when concentrated sulfuric acid is added to chlorates due to traces of ethanol in the pipette I used. I expect the whole beaker of alcohol to just catch fire, or even explosively react with the chloric acid.

So overall, I don't think your method will remove methanol to any significant extent and can even be dangerous.