Sciencemadness Discussion Board

Warning about OTC solvants

trinitrotoluene - 12-8-2004 at 02:07

I always suspected that OTC solvants were not really pure. Indeed this was proven true. I was curious about acetone, so I took a quart (946ml)of acetone and did a distillation on it. The brand that was used was Klean Strip brand, purchased from the Home Depo. I proceeded to distill it using a water bath. This was distilled using a 500ml erlinmyers flasks, some Cu tubing and 125ml RB flask at the receiving end. After 4 hours of distilling, I finally distilled off all the acetone, and I was left with 3ml of this yellowish liquid. It was a good idea I used a water bath, it kept the temperature steady, and once all the acetone distilled away, the other gunk dosen't distill over.

But this yellow liquid has a low viscosity like any other solvant, but when I took a smell it smelled nothing like that ketone acetone. It smelled like it was a hydrocarbon, very simular to the smell of gasoline.

I tossed the last 25ml of acetone that had distilled over, I basically used it to clean the tubing.

As of now, I will distill all my OTC solvants to ensure purity, I suspect other OTC solvants are contaminated with other organic impurties as well. If you need high purity solvants I recommend a careful distillation of those chemicals.

vulture - 12-8-2004 at 04:00

DON'T just distill solvents like acetone and ethers like that! There is always some peroxide content in there which stays behind!

The leftover you saw could very well be a mixture of condensation products/peroxides.

Yes, there is something else in OTC solvents

Democritus of Abdera - 12-8-2004 at 08:42

I had mentioned it several months ago, in this thread Under the username Hermes Trismegistus.

The product you found is a rust-inhibitor, it plagued my younger days when I would use various solvent for the extraction of "herbal" products and made alot of trouble for those of us trying the chlorobutanol synth,

When the acetone/rust inhibitor was oxidized to chloroform, the R.I. must jump right into the chloroform layer, when the acetone/chloroform mixture is reacted to form chlorobutanol, the R.I. must have an affinity for the solid chlorobutanol and lock lips with it.

I found the solution in getting pure and stable white crystals was in pre-distilling both the acetone and chloroform prior to attempting the synth but I'm not certain anyone paid me any mind.

It wouldn't be suprising, because besides being a rank amateur in the chemical arts in comparison to many of the resident illuminati, I'm also very opinionated.

I once argued (very unsuccessfully) with someone on this forum, that the heat capacity of moist air must be significantly different to that of dry air. It turned out I was arguing with a chemical engineer who designed cooling towers for seven years.

Boy, did I feel stupid! Of course being completely wrong never seemed to stop me from getting in an argument though.....:D

If_6_was_9 - 12-8-2004 at 15:29

I've seen something refered to as virgin acetone in stores used for removing finger nail polish. I wonder how pure that is.

If_6_was_9 - 12-8-2004 at 15:31

Check this out:

[Edited on 12-8-2004 by If_6_was_9]

trinitrotoluene - 12-8-2004 at 21:50

Vulture, people tossed metal into diethyl ether to prevent the formation of peroxides. I did the same when I was distilling acetone by tossing silver into there. And I distilled it carefully in a water bath at 80*C.

Saerynide - 12-8-2004 at 22:48

Why does tossng in metal keep peroxides from forming?

kyanite - 13-8-2004 at 04:17

I think it has something to do with breaking down the peroxide bond like using silver in H2O2 engines.
The metal is a catalyst

vulture - 13-8-2004 at 10:50

Catalysing the decomposition of the peroxides does not seem a very good idea to me, as this would only increase the danger of explosion.

The metal is probably there to reduce the peroxides. This can also be done with FeSO4, Na2S, etc.

Democritus of Abdera - 13-8-2004 at 11:03

Originally posted by vulture
Catalysing the decomposition of the peroxides does not seem a very good idea to me, as this would only increase the danger of explosion.

If the action was catalysis, would the regular catalysis of single peroxides pose a real threat?
Would it result in a buildup of gas in the container?
Would the gas layer over the solvent be oxygen enriched?

svm - 13-8-2004 at 11:22

I think the metal does catalyze the decomposition of the peroxides (probably to water). I think it's fairly safe since it keeps the overall amount of peroxides very low. when you get a build up of peroxides and then decompose them all at once, then you have a problem.

so I think that as long as you toss the metal in before you start heating, you should be okay.

unionised - 14-8-2004 at 04:04

If I read your post correctly then in a quart of solvent you found half a teaspoon of impurity.
Are you going to go back to the shop and complain that it's only 99.7% pure?

Just as an aside, do you know what diacetone alcohol and mesityl oxide smell like?
They are formed by boiling acetone in alkaline conditions. Most glass is alkaline.

I wasn't aware that borosilicate was alkaline.

Democritus of Abdera - 14-8-2004 at 05:33

and grateful for the info...

But I can support TNT on one point, the residue. THe apparatus I used was steel and constructed from two steel drums, I'll describe in another thread. (I used a great deal more solvent and also tried various other solvents, acetone, toluene, xylene, methyl hydrate, and most often of all I used white gas (hexane) and the result was always the same, if it came in a metal can, it had the same damn residue (rust inhibitor)

trilobite - 14-8-2004 at 08:32

I wouldn't be afraid of acetone leaving behind an explosive peroxide residue unless the acetone was already used in a way that could lead to formation of peroxides. No solvent handbooks warn of this, someone would surely had an accident if it was as commonplace as with ethers.


Democritus of Abdera - 14-8-2004 at 08:57

the "warning about OTC solvents" that TNT felt worth mentioning was aimed more towards those attempting to obtain a pure product than any possible dangers.

Still, I have read, that as a general rule, it is bad form to distil even close to dryness.

distilling solvents

maj - 18-4-2005 at 08:37

When distilling some solvents (I.E. Ether) Copper is always a welcome addition.

Quince - 18-4-2005 at 22:07

Likely the acetone from a drug store is more pure, and is what I use as a reagent. It is a good deal more expensive, but that's not an issue unless one is doing a large scale synthesis.

I'd like to inquire as to the purity of drain cleaner H2SO4. Someone mentioned that they add buffers to prevent heating in drains, and I've seen a dense impurity in the bottle of the clear drain cleaner. How likely is this to be a problem? Distillation of H2SO4 is a pain.

cyclonite4 - 18-4-2005 at 22:26

Originally posted by maj
When distilling some solvents (I.E. Ether) Copper is always a welcome addition.

Why did you open up an 8 month old thread to mention that? It has already been covered in the Diethyl Ether thread.

Anyways, adding a metal such as copper or silver does not decompose peroxides, it prevents formation of them. If it were a catalyst to peroxide break-down, think how much more trouble it would cause.

Organikum - 19-4-2005 at 07:59

When distilling ether one adds some ferrous sulfate to the flask - this breaks down existing peroxides.
Into the receiving and storage vessel one adds some pices of copperwire what prevents further peroxide formation, does NOT break down existing peroxides though.

I never heard about dangers distilling acetone, I believe this being a myth - at least something what never happens in reallife.


mick - 17-5-2005 at 13:25

Never distil an ether to dryness because the bit that is left could be a peroxide. Di-isopropyl ether is the one to watch, any crystals and be careful.