At some point in our collective chemical hobby experience, we will encounter chemicals that may represent a significant hazard beyond “flammable”,
“corrosive”, or “poison”. Nothing is worse IMHO then explosion hazards, after all…peroxides kill chemists.
The following link is provided for your reading enjoyment. Covered are such topics as: detection, disposal, storage, mechanism of formation, and when
to call the bomb squad.
Google search “peroxide formation” for further information.
--- this has been a public safety announcement ---12AX7 - 26-1-2007 at 14:57
So where does the peroxide come from, anyway? Oxygen most obviously, but (on the other page) they were awfully concerned about that
pressurized cylinder of butadiene -- how the hell does anything peroxidize that?
Timvulture - 26-1-2007 at 15:34
Butadiene will (like any reactive allylic species) tend to form epoxides in prolonged contact with oxygen. Epoxides have a nasty habit to polymerize
violently.12AX7 - 26-1-2007 at 19:33
Okay, but, sealed, positively pressurized, what am I missing here? Oxygen doesn't tunnel through a half inch of steel...garage chemist - 26-1-2007 at 19:45
I agree. The danger was vastly exaggerated and the destruction procedure was absolute overkill. But thats probably how it goes today, after all,
exaggeration of dangers and unnecessary procedures create jobs. The downsides of that movement today are only too well known to us (restriction of
chemicals).
I looked up the "popcorn polymer" that does form in unstabilized butadiene (also in absence of oxygen) and its dangers are that it is spontaneously
flammable in air (but not explosive) and may block butadiene pipes leading to rupture because of pressure buildup. Nothing of concern in the anoxic
butadiene cylinder.
Solvents like Ether and Tetrahydrofuran are what one has to worry about.
These two are however always delivered with stabilizer today and remain peroxide-free for at least a year.
My 2-year-old 1L bottle of Ether is still peroxide-free (tested with KI solution) despite frequent opening.
However, both drying and distillation remove the stabilizer and makes them become dangerous. This must be kept in mind- either use them up quickly, or
restabilize.
[Edited on 27-1-2007 by garage chemist]Dr.3vil - 26-1-2007 at 21:02
I think what GC mentioned regarding THF and ether is extremely relevant to amateur scientists, especially when attempting to render such solvents from
OTC sources. THF is commonly found in many brands of PVC pipe cement in concentrations of up to 40%. Ether from starting fluid often contains an
aliphatic hydrocarbon or other additives.
Separating desired chemicals by distillation is simple enough, however increased heat and exposure to atmospheric oxygen and light present a challenge
to safely recovering the product. Perhaps I should continue this thread in reagents and acquisition as ultimately I would like to establish a best
practice for recovery of THF and ether from OTC sources. Certainly not distilling to dryness and adding a stabilizer to the recovery flask would seem
prudent.
How oxygen got into the cylinder, I think we will never know. Never underestimate the incompetence or carelessness of industrial operators. Bhopal
anyone? Chernobyl?Levi - 27-1-2007 at 02:42
Quote:
Originally posted by 12AX7
Okay, but, sealed, positively pressurized, what am I missing here? Oxygen doesn't tunnel through a half inch of steel...
Sure it does! Though it leaves a trail of rust behind it and takes a looong time...quicksilver - 27-1-2007 at 07:54
"when to call the bomb squad" is never unless you want to be prosecuted or persecuted IMO. In this age we live in the responsible thing to do may be
one's undoing. I deeply believe that if one feels that an experiment may be uncontrolable to the point of needing outside intervention, one should not
undertake that experiment. The press is insatiable in it's feeding frenzy on "drug labs", etc. The neighbors would be unmerciful. That is not to say
that experiments with energetic materials or toxics should be avoided but if one is awair of one's limitations - the worst will always be managable. I
think it's important that we be as cautious as possible that the repercussions of our hobby or work have upon the continuation of private
experimentaion. We can easily promote or destroy chemistry as a hobby. I certainly don't mean to sound harsh in this but I think that if (as an
example) a hobbiest work with chemistry is not cognisant of the impact he has upon the public's perception of that hobby....it will quickly be
unavailable for further experiments. The press can be damning.Dr.3vil - 27-1-2007 at 08:13
Quote:
Originally posted by quicksilver
"when to call the bomb squad" is never unless you want to be prosecuted or persecuted IMO. In this age we live in the responsible thing to do may be
one's undoing. I deeply believe that if one feels that an experiment may be uncontrolable to the point of needing outside intervention, one should not
undertake that experiment.
I have a 30 year old brown glass bottle of amyl alcohol. Is it safe to open it?
As far as I know amyl alcohol does not form peroxides but I could be wrong.
[Edited on 11-2-2007 by Zinc]joeflsts - 18-2-2007 at 17:05
Another good website for peroxide information. Also @Zinc I'm not aware of Amyl Alcohol forming peroxides.
I agree. I'm just aware of amyl alcohol forming nauseous moments if you accidentally smell it. thefips - 24-2-2007 at 12:13
Is the diethyl ether peroxide soluable in the diethylether?
Or could it appear, that crystals can fall out and settle on the bottom of the bottle? I have a bottle with diethyl ether and there are small
crystals on the bottom, but the peroxide test was negative.
Has someone an idea what it could be, or was my peroxide test wrong? I used KI in 60% acetic acid with half the volume ether in a test tube with
shaking.Magpie - 24-2-2007 at 12:33
Here is a test for peroxides in ether from my old lab manual:
Add 1 mL of a 10% solution of KI to 10 mL of water in a small test tube and acidify the solution with a few drops of dilute sulfuric acid. Add 2mL
of ether and shake the tube for a moment. Add a few drops of starch paste indicator or a strip of starch-iodide paper and shake. Appearance of a
blue color indicates the presence of peroxides in ether.joeflsts - 25-2-2007 at 11:16
Another test: 10% by weight potassium iodide (KI) mixed with dH2O. About 1ml of this with 10ml of your test material, mixed, shaken after a 30 minute
rest will display a color change if peroxides are present.
The peroxide is oxidizing the KI kicking out I2. The I2 gives a purple / brownish solution depending on the solvent.
Purple, brown or purple brownish indicate a high conc. of peroxides. Yellow indicates a lower concentration.garage chemist - 25-2-2007 at 12:05
The peroxide is soluble in the ether, yes. It does never precipitate and with diethyl ether, only becomes dangerous during evaporation of the ether.
Your peroxide test sounds good, it should have detected any peroxide.
The crystals on the bottom might be a drying agent (molecular sieves, KOH, CaCl2 etc...) or maybe some sort of stabilizer. Peroxides are the only
thing they surely aren't.
Spiked my curiousty
joeflsts - 25-2-2007 at 13:25
About 4 months I ago I made some Ethyl Acetate. Today I decided to do a peroxide test on a few items:
3% H2O2 - Tested Positive
35% H2O2 - Test positive
Tap Water - Negative
Ethyl Acetate - Test positive with low concentration (yellow)
null
It does not appear that Ethyl Acetate is a known peroxide forming compound. I'm curious - is it possible I'm seeing another reaction going on here
that is liberating I2?
Joethefips - 25-2-2007 at 13:54
Thanks a lot for the answers. Toworrow I will test the diethyl ether again, because it is standing around for about 3 years now.
And there is a bottle of THF, too, which I should test.
I think it is a very important thing to know if the chemicals are as save as they should be.pantone159 - 25-2-2007 at 14:52
Quote:
Originally posted by joeflsts
It does not appear that Ethyl Acetate is a known peroxide forming compound.
* Prepare reagent by adding 100 mg sodium iodide (NaI) or potassium iodide (KI) crystals to 1.0 ml of glacial acetic acid. Add 0.5 to 1.0 ml of
material being tested to an equal volume of reagent. A yellow color indicates a low concentration (~0.1 per cent) and brown a high concentration of
peroxide in the sample. A blank should be run, using some non-peroxidizable compound such as pure n-hexane.
* Peroxide test strips, which change color to indicate the presence of peroxides, may be purchased through most laboratory reagent distributors.
For proper operation, the strips must be air-dried until the solvent evaporates and then exposed to moisture.
There are two effective methods for removing peroxides safely:
1. Pass the solvent through a short column of activated alumina. No water is thereby introduced. The alumina catalyzes the decomposition of many
peroxides, but it is possible that some peroxide will be retained unchanged on the column. The alumina should therefore be disposed of as a flammable
material.
2. Make up a reducing solution from 60 g ferrous sulfate (FeSO4), 6 ml concentrated sulfuric acid (H2SO4), and 110 ml water. Shake the sample with
this solution to remove the peroxide.
From H.L. Jackson, W.B. McCormack, C.S. Rondestvedt, K.C. Smeltz, and I.E. Viele: "Safety in the Chemical Laboratory LXI: Control of Peroxidizable
Compounds". J. Chem. Educ. 47(3): A176 (March, 1970).Filemon - 24-5-2007 at 12:15
Would it also work FeCl3 to remove the peroxides?Mumbles - 24-5-2007 at 13:26
Maybe.
The mechanism by which it removes the peroxide is to reduce the bond. It reduces the peroxide, and it gets oxidised to Fe 3+ The only reason I say
maybe is that transition metals are known catalysts to decompose peroxide bonds. I wouldn't count on it to get rid of peroxide contamination in your
solvents and such.12AX7 - 24-5-2007 at 15:46
Incidentially, I noticed iron undergoes a reaction, in strong acid at least (>>1M sulfuric). The H2O2 dissipates quite quickly. It was a
purple color, I think. It's been a while since I did it.
