RxnJackson
Harmless
Posts: 23
Registered: 3-11-2007
Location: The mean streets of Rage
Member Is Offline
Mood: No Mood
|
|
Storage of Di-ethylether
I understand that anhydrous ether produces explosive peroxides during storage. Is there anyway to determine if a tin of ether needs to be replaced
(OSHA reccomends six months, however I'm wondering if there is a practicle way to determine this, thus not having to throw out and re-order chems that
are perfectly fine, and expensive I may add).
This generation of peroxides is an oxidation from air or H2O? Is there a prefered storage container to reduce this?
Last question, the peroxide generated, how does it cause an explosion and can it occur with no ignition source?
RxnJackson
EDIT: If this topic belongs in the organic chemistry section, could a moderator please move it. I wasn't sure if this forum was a reagent discussion,
or just reagent aquisition discussion.
[Edited on 4-11-2007 by RxnJackson]
|
|
|
smuv
National Hazard
Posts: 842
Registered: 2-5-2007
Member Is Offline
Mood: Jingoistic
|
|
I would search the forum about this one, I am sure it has been discussed quite a bit. Also, check the forum library and download "Practical organic
chemistry" by Vogel, it has a pretty detailed discussion about ether peroxides, and removing them.
Yes peroxides can cause explosions without ignition sources, usually peroxides are only a problem when they begin crystallizing out of the ether (in
solution they are less sensitive). Ether peroxides are shock sensitive.
|
|
|
feacetech
Hazard to Others
Posts: 163
Registered: 12-2-2007
Member Is Offline
Mood: No Mood
|
|
Organic peroxides will form in many organics over time if exposed to light or catalyst and air.
store in the dark/amber bottle to help reduce this.
orgainc peroxides are very unstable since they contain oxidser and fuel in one molecule. With sufficent contaimination even the slightest vibration
can set one off.
from wiki:Consequently it is recommended that ether be stored over potassium hydroxide, which not only destroys peroxides but also acts as a powerful
desiccator (drying agent). Extreme care must be taken with samples showing signs of crystal growth or precipitates.
|
|
|
RxnJackson
Harmless
Posts: 23
Registered: 3-11-2007
Location: The mean streets of Rage
Member Is Offline
Mood: No Mood
|
|
Thank you all, I am very farmiliar with Vogel's book and have a copy here, I surely should have reviewed that in detail.
My current storage system is a Fisher chemical Tin for anhydrous liquids, should be sufficient.
The crystal growth would appear as a white crystialline, correct? Anywhere there is an air/solvent boundry, or within the container itself? Just
curious as to what I should be looking for in the future.
|
|
|
jam640
Harmless
Posts: 29
Registered: 3-10-2007
Location: Rainbow Island
Member Is Offline
Mood: Fat
|
|
I remember when I started studying chemistry how scared I was of using diethyl ether due to the peroxide buildup stories we were told in lectures. In
real life most commercial diethyl ether is stabilized by some additive such as BHT, and peroxide formation is really not an issue unless you plan to
store your solvent for a very long time. The maximum recommended time stated on my bottles of ether from Fisher is 2 year after opening. I have never
experienced peroxides and I've used bottles that was opened over a year ago. But as 1L bottles of diethyl ether are pretty cheap (atleast where I
live, guess it varies but still when it comes to safety it's worth to cash out a few bucks more), why not just buy the amount you need at the moment
and use it. A liter of ether is used up pretty fast when working with any organic chemistry. 
[Edited on 5-11-2007 by jam640]
|
|
|
RxnJackson
Harmless
Posts: 23
Registered: 3-11-2007
Location: The mean streets of Rage
Member Is Offline
Mood: No Mood
|
|
quote: "why not just buy the amount you need at the moment and use it. A liter of ether is used up pretty fast when working with any organic
chemistry."
Well said....
|
|
|
smuv
National Hazard
Posts: 842
Registered: 2-5-2007
Member Is Offline
Mood: Jingoistic
|
|
Well...If you are recycling your solvents (which I think should be done for ecologic and monetary reasons), distilling ether will remove the bht
inhibitor. With fresh ether yes you can rely on the inhibitor, but when recycling solvents the inhibitor is lost (assuming you are purifying w/
distillation). Also when boiling/evaporating ether to dryness I think it is a good idea to check for peroxides before hand, even if it is inhibited.
|
|
|
garage chemist
chemical wizard
Posts: 1803
Registered: 16-8-2004
Location: Germany
Member Is Offline
Mood: No Mood
|
|
My 2 year old and often used 1L bottle of ether (brown glass, and a copper wire spiral at the bottom since I read copper inhibits peroxide formation-
no idea if the copper is having any real effect) is still entirely free of peroxides.
As a general rule, stabilized ether in its original bottle is very safe for a long time.
The problems start once you dry and distill it for use in e.g. Grignards- store such unstabilized ether airtight and only for very short terms. Best
dry only the amount of ether you will be using immediately for the reaction on this day, and restabilize any excess dried ether.
In unstabilized ether peroxide formation is indeed rapid (peroxide test clearly positive after 2-3 days) and a very real danger if it is to be
distilled.
|
|
|
RxnJackson
Harmless
Posts: 23
Registered: 3-11-2007
Location: The mean streets of Rage
Member Is Offline
Mood: No Mood
|
|
good to know. Looks like I can store the material (as it is stabilized, I checked the bottle and contacted Fischer today), and dry material as I need
it (and in only the amounts I will be using).
I probably will not be recycling solvents (though it is certainly more eco-friendly), but distilling only neccessary material for a gringard, or
wittig as needed makes good sense.
|
|
|
smuv
National Hazard
Posts: 842
Registered: 2-5-2007
Member Is Offline
Mood: Jingoistic
|
|
@ GC I too store my ether with copper wire in it, as I have read the same about copper and iron (i think it was iron) destroying peroxides. I have
noticed that if you take a somewhat dull piece of copper wire, and store it in ether for a few days it brightens visibly. Maybe this means it works?
I would assume copper inhibits peroxides as they oxidize the Cu(0) to Cu(I) or Cu(II).
One of these days I should run a long term experiment to see how well the Cu spiral method actually works (Actually I will find out very soon as 1.5
months ago i stupidly forgot to inhibit about 250ml of freshly distilled ether, and simply stored it over a copper spiral). Bht is readily available
OTC (in my country) as it is supposidly an alternative treatment for herpes...
|
|
|
evil_lurker
National Hazard
Posts: 767
Registered: 12-3-2005
Location: United States of Elbonia
Member Is Offline
Mood: On the wagon again.
|
|
For ether to form peroxides, it needs water, oxygen, and a radical initiator, that being light or certain metals... IMO light being the worst
offender.
As such, one of the best peroxide stabilizers is plain old 4a molecular sieves... given enough time they will actually remove the hydro peroxides in
ether. BHT is good too, and it only takes like .001% IIRC.
Aluminum bottles are preferable since they keep the ether pitch dark, and are less likely to have vapor pressure issues, thus they are used in
commerce for shipment. Amber glass bottles are #2 and have the benefit of being able to visibly detect peroxides crystals around the cap. Deep
freezers are also very nice for ether storage, but present a remote explosion hazard unless they are certified for lab use.
Not all chemicals are bad. Without chemicals such as hydrogen and oxygen, for example, there would be no way to make water, a vital ingredient in
beer.
|
|
|
Centimeter
Harmless
Posts: 28
Registered: 19-5-2004
Member Is Offline
Mood: No Mood
|
|
You asked how these explosive properties are evolved so I thought that I would explain that, as it is a rather fascinating mechanism. The " * "
denotes a single electron.
O=O ---> *O-O*
*O-O* + CH3CH2-O-CH2-CH3 ---> CH3-CH2-O-C*H-CH3
CH3-CH2-O-C*H-CH3 + *O-O* ---> CH3-CH2-O-CH-CH3
------------------------------------------------------------------l
-------------------------------------------------------------*O-O
-The slightest bump can cause the attached oxygen radical to swing up and snatch one of the hydrogens off of the second carbon. This process is not
all that exothermic. However it results in some e- rearrangements...
CH3CH-O-CH-CH3 ---> 2 CH3CH=O + *OH
-------*-------l
---------------O-OH
*OH + ( )-H ---> H2O + lots of heat
The bond dissociation energy of H-OH is 119 kcal/mol. Thus, it'll snatch up any hydrogen it can and in quite an exothermic manner. My professor refers
to hydroxide radicals as Blackwater gunmen.
I hope that this is correct!
Edit- Holly crap is this hard to make the text images work!
[Edited on by Centimeter]
[Edited on by Centimeter]
[Edited on by Centimeter]
[Edited on by Centimeter]
[Edited on by Centimeter]
|
|
|
RxnJackson
Harmless
Posts: 23
Registered: 3-11-2007
Location: The mean streets of Rage
Member Is Offline
Mood: No Mood
|
|
Interesting centi. I like evils concept about the 4A sieves. These would be effective in maintaining dryness as well and might eliminate the
necessity to re-distill the material for anhydrous reactions (ie: butylLI or other STRONG bases).
|
|
|
smuv
National Hazard
Posts: 842
Registered: 2-5-2007
Member Is Offline
Mood: Jingoistic
|
|
@ evil_lurker, I have never heard that bit about water being necessary to form peroxides. I have found a source that shows water in ether, can form
peroxides via one mechanism, but shows there is also another mechanism, in which oxygen alone can produce peroxides. I was wondering if you could give
me a reference for your statement, as I would really really like to believe that keeping ether dry inhibits peroxide formation, but at least one
source (albeit dated, from the 50's at least) says that this is not the case.
See the attached pdf
[Edited on 5-11-2007 by smuv]
Attachment: ether_peroxides.pdf (59kB)
This file has been downloaded 951 times
|
|
|
Centimeter
Harmless
Posts: 28
Registered: 19-5-2004
Member Is Offline
Mood: No Mood
|
|
smuv, as you can see from the mechanism that I posted, only oxygen is neccesary to form the explosive product. Keeping your ether dry will have no
effect on this fact. Keeping it away from air or providing a means to neutralize oxygen radicals, however, can potentially prevent explosive
byproducts from forming.
IMHO, situations dealing with explosives, especially peroxides, ought to be avoided at all cost. I've heard of bomb squads being called in to deal
with diethyl ether that's sat for too long. Hence I would highly recomend using a different solvent or using only fresh ether.
[Edited on by Centimeter]
|
|
|
RxnJackson
Harmless
Posts: 23
Registered: 3-11-2007
Location: The mean streets of Rage
Member Is Offline
Mood: No Mood
|
|
| Quote: | Originally posted by Centimeter
IMHO, situations dealing with explosives, especially peroxides, ought to be avoided at all cost. I've heard of bomb squads being called in to deal
with diethyl ether that's sat for too long. Hence I would highly recomend using a different solvent or using only fresh ether.
[Edited on by Centimeter] |
or stabilized ether within it's expiration, correct?
|
|
|
Centimeter
Harmless
Posts: 28
Registered: 19-5-2004
Member Is Offline
Mood: No Mood
|
|
Worry not, I bring an old thread up with good reason!
I found an interesting section in the book, Introduction to Organic Chemistry by Andrew Streitwieser that pertains to this subject.
"A simple test for peroxides is to shake a small volume of the ether with aqueous KI solution. If peroxides are present, they oxidize I- to I2. The
iodine complexes with excess iodide to form triiodide I3-. The characteristic brown color of I3- is diagnostic for the presence of peroxides...
Contaminated ether may be purified by shaking with aqueous ferrous sulfate to reduce the peroxides."
The book says some other important things about the issue. It recommends using methyl t-butyl ether as it is not as prone to forming peroxides as
ethyl ether. It also says that distilling questionable ether to dryness (i.e. anything that you make at home!) is not a good idea as the peroxides
will concentrate in the boiling flask and pose a large explosion hazard.
|
|
|
smuv
National Hazard
Posts: 842
Registered: 2-5-2007
Member Is Offline
Mood: Jingoistic
|
|
I hate to say, but while good info, this is all very standard stuff. More methods of detection and removing peroxides are outlined in "Practical
organic chemistry" by Vogel which was referenced above.
|
|
|
guy
National Hazard
Posts: 982
Registered: 14-4-2004
Location: California, USA
Member Is Offline
Mood: Catalytic!
|
|
| Quote: | Originally posted by Centimeter
You asked how these explosive properties are evolved so I thought that I would explain that, as it is a rather fascinating mechanism. The " * "
denotes a single electron.
O=O ---> *O-O*
*O-O* + CH3CH2-O-CH2-CH3 ---> CH3-CH2-O-C*H-CH3
CH3-CH2-O-C*H-CH3 + *O-O* ---> CH3-CH2-O-CH-CH3
------------------------------------------------------------------l
-------------------------------------------------------------*O-O
-The slightest bump can cause the attached oxygen radical to swing up and snatch one of the hydrogens off of the second carbon. This process is not
all that exothermic. However it results in some e- rearrangements...
CH3CH-O-CH-CH3 ---> 2 CH3CH=O + *OH
-------*-------l
---------------O-OH
*OH + ( )-H ---> H2O + lots of heat
The bond dissociation energy of H-OH is 119 kcal/mol. Thus, it'll snatch up any hydrogen it can and in quite an exothermic manner. My professor refers
to hydroxide radicals as Blackwater gunmen.
I hope that this is correct!
Edit- Holly crap is this hard to make the text images work!
[Edited on by Centimeter]
[Edited on by Centimeter]
[Edited on by Centimeter]
[Edited on by Centimeter]
[Edited on by Centimeter] |
The triplet state of oxygen (*O-O*) is the most stable form of oxygen and does not require any light excitation. The singlet state (O=O) is unstable.
The formation of a CH* radical is formed by light breaking the C-H bond. Since triplet oxygen can react with the doublet state of the CH* radical,
it will form a RO-HC-OO* radical which will react with the H* radical to form RO-HCOOH
|
|
|
![[*]](images/xpblue/default_icon.gif){kind=icon}
