When I was at work today there were some people who placed some PVC pipe. They used special glue called SABA. I read the contents and the line “can form explosive peroxides” stroke my eye :D It contained tetrahydrofuran, it burns very fast with a crackling effect.
I searched and found this:

Tetrahydrofuran
Synonyms: THF; 1,4-Epoxybutane; Butylene oxide; Cyclotetramethylene; tetramethylene oxide; oxacyclopentane; Cyclotetramethylene oxide; Furanidine; Hydrofuran; oxolane;

Registry number: 109-99-9
Density: 0.89
Melting point: -108 °C
Boiling point: 66 °C
nD20: 1.407
Flash point: -21 °C
<img src="http://www-woc.sci.kun.nl/data/dadml/2d/gif/109-99-9.gif" alt="" />
Stability
Stable. Incompatible with halogens, strong oxidising agents, strong reducing agents, strong bases, oxygen. May generate explosive peroxides in storage if in contact with air. Highly flammable. Store at room temperature under nitrogen. Hazardous polymerisation may occur. Light sensitive. May contain 2,6-di-tertbutyl-4-methylphenol (BHT) as a stabiliser.

Hazard Symbol
- F Highly flammable
- Xi Irritant

Risk Description
- R11 Highly flammable.
- R19 May form explosive peroxides.
- R36/37 Irritating to eyes and respiratory system.

Safety Description
- S16 Keep away from sources of ignition - No smoking.
- S29 Do not empty into drains.
- S33 Take precautionary measures against static discharges.

Can someone tell me wich explosive peroxides are being generated?

maybe:
-C-O-C- => -C-O-O-C-

(couldnt make a ASCII draw)

not exactly as pretty ar your picture, but a carbon-ring including a peroxide bond is very "tense", thus unstable..so by introducing a second oxygen atom you create the peroxide

/rickard

<small>[ May 22, 2002, 03:01 PM: Message edited by: rikkitikkitavi ]</small>

I`ve heard of THF. It`s used as a solvent for PVC (like you said), and lots of other stuff. I have a bottle of a very thick liquid used to "weld" plastic. It contains:

30-60% Methylethylketone [as in MEK-peroxide :) ]
10-30% Toluene
10-30% THF

It doesn`t say anything about "...may form explosive peroxides" though.

Maybe it can be used to make an "interesting" mixture of MEKP and that "explosive peroxide" ...but it`ll probably be unstable, extremely toxic, or other stuff that keep us from using it :(

Are there any ways of extracting the THF?
What is hazardous polymerisation?

THF BP 66 C, miscible with water in all proportions

MEK BP 78,6 C, solubility 24 %in water

Toluene BP 111 C, inmiscible with water

=> mixing with water would separate the toluene from the others,
better would be to first distill of THF and MEK from the toluene, and then adding a little water to the distillate, just engough to dissolve the THF, giving two liquid phases ,water-THF and MEK

/rickard

The THF may undergo a polymerisation reaction forming long chains (presumably between the lone pairs on the oxygen and the pi bonds on the ring), liberating heat and possibly leading to thermal runaway.

IIRC THF peroxide does contain -C-O-O-C- in the ring.

Have you seen my site? If so, look at the huge great mushroom cloud thingy on the homepage. That's 15 gallons of partially peroxidised THF going off.

I've often wondered what that was. :) Awfully impressive. What was the source of the image? Was this an intentional disposal of some THF that had been improperly stored in the presence of oxygen, or what?

It came from the website of a company that specialised in disposing of hazardous chemicals, including organic peroxides. I rather liked the look of it so I stole it for my website.

How can my THF be used to generate THF-peroxides?

P.S. I found this statement on <a href="http://www.genevac.co.uk/" target="_blank">Genevac</a>
</font><blockquote><font size="1" face="Verdana, Arial, Helvetica">quote:</font><hr /><font size="2" face="Verdana, Arial, Helvetica"> Peroxide Explosions
There are reports in the literature of peroxide explosions arising from distillations of THF. Such explosions do not need the presence of air and can happen spontaneously if there is too high a peroxide content in the THF.
</font><hr /></blockquote><font size="2" face="Verdana, Arial, Helvetica">

The explosions themselves do not require air, if peroxides are already present in appreciable amounts. The formation of the peroxides in the first place, however, does require a source of oxygen radicals.

Yikes! The purposeful preperation of THF peroxides strikes me as something you do not want to do. If you wish to create chaos in your life, just set a bottle of THF out with a bubbler bubbling air into it. Try a fish pump with an airstone at the end. Assuming your fishstone dosn't dissolve (mine always seem to in solvents, sigh) you will get some peroxide. The longer you let it sit, the more you will get. For added preservation of your THF try this in a reflux setup.

Do not try to isolate the solid product. You would be advised to place this in a vacuum or use an aspirator to let the solvent evaporate down some. The crystals respond to heat, shock, and even negative thoughts. Don't say I didn't tell you so when you get killed doing this. Of course knowledge is power.

Hmm...
What if a stupid and mis-informed individual was to have about a 20gallon drum of THF, placed it in a vacant field, added a couple gallons of 50% H2O2 and a bit of HCl.... what would be the result?

Im guessing you wouldnt want to be NEAR the bitch when it was mixed... but what would -theoretically- happen? Would it simmer a bit and then detonate? Would it require to evaporate first?

Oh, Mr Cool, when you say "Going Off" what do you mean? Was a base charge used? A small blasting cap? Was it... poked?

Hmmmmmmmm....
*ponders*

My guess that adding peroxide and hydrochloric acid to THF would do absolutely nothing, at least not what you expect. Remember that THF is an ether, not a ketone like acetone. Ethers are very unreactive to just about everything. With hydrochloric acid and some heating you would most likely end up with some dichlorobutane and a bit of butanediol, as ethers undergo cleavage with HX acids. However, even this is unlikely as HCl is not the reagent of choice to cleave an ether.

Ethers are very reactive with oxygen gas. Bear in mind that the peroxides of ethers are not attached to the ether oxygen. There will be an –OOH peroxide attached to one of the adjacent carbons as a side chain. This is an ether peroxide, not just a peroxide. These are called hydroperoxides. If you just set your THF out in the open air, you will get peroxides, and they will be dangerous, and they will form other nasty peroxides.

<small>[ June 27, 2002, 09:54 PM: Message edited by: megalomania ]</small>

What about diazomethane, dinitrogen dioxide, liquid ozone, 90%+ H2O2, or diazoborane as oxygenators?

Two components mixed then shocked. Would they not explode? Or at least burn with incredible intensity like rocket fuel.

Err, when using ozone or 90%+ H2O2 you won't have to shock it anymore, let alone you'll still be able to shock anything...

and you would get liquid O3 from... _where_ ????
<img border="0" title="" alt="[Eek!]" src="eek.gif" />

That shits like THE most evil gas on the planet..... kills anythin and everything nice 'n fast. Even more fun than Hydrogen fluoride..eheheh

The only genuine use of ozone I can think of is in large scale water purification systems... as it kills any and all weeblie's that swim in tap water :mad:

Would industrial gas supply stock it? Or would one make it ones-self... HV thru liquid oxygen - ALSO somewhat 'tricky'

Oh, and strong ass H2O2 -HTP- was\is used in torpedoes as a portable oxygen source.
Its really frikkin dangerous (fun) tho...

That stuffs measured in 'Volume' = how many times the stuff increases in volume when decomposed.
Eg: 30% (or there abouts) is 100volume.
Thus 60% would be 200vol... and 90% would produce a whopping 300x its volume when decomposed.
(are these figures correct? im rather tired...)

Think of the fun you could have with that.

Oh, and the fact that it dissolves pretty much any organic matter in a puff of oxygen :D

If you can get KMnO4, H2SO4 98%, liquid N2 and ozone resistant tubing and reactor vessels you can produce liquid O3.

2KMnO4 + H2SO4 -&gt; K2SO4 + Mn2O7 + H2O

heat
Mn2O7 -&gt; 2MnO2 + O3

Ozone is either made through electrolysis of cold(-50-60C) percloric acid at high current densities or through "dark discharge" (look at USPTO for patents) of high voltages.

Ozone has some use as a environmently friendly bleaching agent in the pulp industry and as water purification agent. It is always made on-site and never stored due to its high toxcitiy and reactivy.

/rickard

Ozone is also generated easily by silent static electrical discharge passed through a stream of oxygen.

3O2 --&gt; 2O3

See <a href="http://www.orgsyn.org/orgsyn/prep.asp?prep=cv3p0673" target="_blank">http://www.orgsyn.org/orgsyn/prep.asp?prep=cv3p0673</a> for details on producing a laboratory ozonizer apparatus.

Ozone is used in organic analysis for ozonolysis of alkenes, forming aldehydes and ketones upon workup, or if extreme oxidizing or reducing condition are used, can form acids or alcohols. This is currently being discussed as an avenue for practical application to synthesis instead of analysis, in another topic at another forum.

As for THF, I'm fairly certain that the peroxide structure would be of coordinate valence complexing on the ether oxygen, looking like compound A rather than the previously proposed structure looking like compound B.

<img src="http://www.methinfo.com/boards/general/binaries/18/18544.gif" alt=" - " />

THF peroxides are extremely dangerous and unpredictable, and should be avoided. Their formation is catalyzed by free radical promoters, and is itself a radical facilitator, as most peroxides are. THF should preferably be stored under slight reducing conditions to inhibit peroxide formation, such as under a nitrogen or noble gas atmosphere, with a small quantity of NaHSO3 or sodium dithionite.

Bee careful please.

PrimoPyro

Primepyro, in structure A, you've got an oxygen molecule with four bonds, but I see no other changes in charge?

EDIT: If you mean a complex, how could the two oxygen molecules attract eachother? They have to be charged then, one negative and the other positive and oxygen with a positive charge can only exist very briefly as a radical in certain reactions due to it's high electronegativity of 3,5.

<small>[ August 06, 2002, 04:59 AM: Message edited by: vulture ]</small>

I just didn't draw it. The central oxygen would be positively polarized I suppose. Technically it doesnt have to be polarized you know.

Look at the structure of oxyacids like chloric acid and chlorates or nitric acid, in fact ALL oxyacids "seemingly" excced the octet, but they don't. It's a basic principle of coordinate valency that is present in ALL oxyacids and I am supposing in this peroxide structure as well.

The only reason I propose this structure is because the other structure would need to open the ring in order to form, and this is very unlikely to happen in an ether. Even if it did I doubt it would re-close so cleanly, I'd suspect polymerization involving a repeating -O-(CH2)4- function.

I can explain the prnciple of coordinate covalent bonds if need be. But I won't force it on anyone. (A lot of people simply don't care I guess.)

PrimoPyro

In oxyacids like HClO3 or HNO3 Cl and N are the central atom and they have a "possible" oxydation state there.

Whereas in your molecule the O is the central atom.
Besides, diethyl ether also forms explosives peroxides with ring opening.

EDIT: Besides, the structure you drew there is not a peroxide, peroxides are always R-O-O-R

<small>[ August 08, 2002, 11:28 AM: Message edited by: vulture ]</small>

The formation of ether peroxides is a radical mechanism where an O2 radical .O-O. attacks the C next the O in the ether R-O-(CH2)-R, forming an ether radical R-O-(.CH)-R and a hydroperoxide radical
.O-O-H , these combine to form an ether hydroperoxide, which I tried to draw with ascii but failed miserably, just think of the C with the R-O-, H-, -R and -O-O-H attached.

Another .O-O. attacks, abstracts an H from the hydroperoxide, so the
-O-O-H attached to the C changes to -O-O. while you get another hydroperoxide radical .O-O-H floating in the solution.

Your etherhydroperoxide radical with the C-O-O. combines with an ether radical R-O-(.CH)-R, so in the end you get 2 ether molecules, where the C next to the O is connected with an -O-O- peroxide bridge.

These pages are the best I could find at the moment explaining radical reactions:
<a href="http://www-users.york.ac.uk/~chem77/Tale_of_2_Radicals_Part_2.htm" target="_blank">http://www-users.york.ac.uk/~chem77/Tale_of_2_Radicals_Part_2.htm</a>
<a href="http://www.usm.maine.edu/~newton/Chy251_253/Lectures/Free%20Radicals/FreeRadicalFS.html" target="_blank">http://www.usm.maine.edu/~newton/Chy251_253/Lectures/Free%20Radicals/FreeRadicalFS.html</a>

There is no such thing as four valence oxygen, which was shown in diagram A.

The valence of an atom is only "extended" from the octet rule by atoms more electronegative being bonded to it - for example HNO3, HClO4, HSO3F, and so on... also, many of the bonds in those substances are ionic.

I already stated that the drawing was simply incomplete. I admit to being incorrect regarding peroxyl structure, but I can assure you that these bonds exist. They are often seen in heteroatomic bonding (as in the atoms bonded are not the same atom) but I extended the possibility to include a homoatomic bond here.

These bonds are derived from coordinate covalent complexes in which both atoms are of strongly (some exceptions like sulfur and phosphorus) electronegative character. The central atom, already achieving its octet through conventional bonding, desires no more bonds. But these other atoms, often oxygen, chlorine, and fluorine are able to attach anyways, seemingly breaking the simple octet rule. But this is not so.

In the compound H2SO4 for example, the structure looks like H-O-S-O-H with an additional two =O groups on the sulfur atom. Being in the same group as oxygen, sulfur's electron configuration is the same, but of one higher energy level. It should be sustained easily as H-O-S-O-H and nothing else, but it is not. This compound does not exist except very weakly possibly in solution. The other two oxygens are able to bond to the sulfur atom with a double bond each, seemingly exceeding it's octet. The sulfur is NOT positively polarized by the way.

Each of the oxygens in the =O groups take two electrons from the sulfur atom to complete their octet. This forms the covalent bond. The electrons taken are derived from the full s orbital in the valence shell of the sulfur atom (two electrons) and the single full p orbital in the valence shell (2 more electrons) the other two p orbitals are both half full, and are busy coordinating the single bonds with the OH groups.

You may ask, "But it takes two electrons to form a covalent bond, so if only two electrons are given from the sulfur to each oxygen, how can they be of double bond strength and configuration?"

The answer is simple: Because the electrons themselves mean nothing. They are simply glue that tie together orbitals. Orbitals are what matter. There may be only two electrons involved in the bond, but they are in different locations. The electrons coming from each of the single S orbitals end up coordinating to TWO seperate p orbitals (each already half full) in the =O atom. This caused both p orbitals to overlap with the sulfur orbital it is involved with, yielding a double bond.

These oxygenated compounds of sulfur are much much more stable than their unoxygenated counterparts. The stability is derived from the multiple electron resonance hybrids able to form with increasing electronegativity of the O=S=O group. The OSO group is very fucking stable and very fucking electronegative to put it bluntly.

This is why it is used in all mega leaving groups like tosyl brosyl nosyl mesyl and triflyl (best leaving group in existence) because they are so stable due to resonance. I bet if I looked hard enough through Beilstein, I could find examples of hypervalent oxygen-oxygen bonds.

PrimoPyro

<small>[ August 08, 2002, 03:00 PM: Message edited by: PrimoPyro ]</small>

The bonding in the sulfate ion is *not* covalent - it is semipolar (combination of ionic and covalent). The sulfur atom theoretically has a charge of +2, the oxygens all have a charge of -1, for an overall charge of -2 for the sulfate ion. All of the oxygens are bonded to the sulfur in the same manner in the sulfate ion. The reason I said "theoretically" is because in reality the values are slightly different; for example, the charge of the sulfur atom is about +1.77 in reality (because in reality, bonding isn't so exact as it is on paper). The s orbital and all three p orbitals take part in the bonding. It's not surprising that the sulfate ion would be semipolar, because the addition of one electron to an oxygen atom is exothermic.

Sulfate ION is polar/ionic because the protons have shifted off into solution. The charge bearing particles are the oxygens. The sulfate ion contains a less electrically influenced sulfur than a neutral sulfate group.

As the =O groups draw electrons away from the sulfur, in the sulfate ion, the O- groups have a very strong electron pushing effect and neutralize the S group entirely, sometimes too much, which of course is what leads to the resonance of the electron throughout the molecule, reducing one =O to -O- (-O negative) and the previously -O- becomes a =O. This only happens in the ion when there is free charge. This is why the ion is so stable.

PrimoPyro

I just realized that we're debating about which of two accepted theories is correct. Look towards the end of this page, at "3. The sulfate ion in terms of hybridization and resonance theory." You were backing the first theory, I was backing the second.

<a href="http://www.madsci.org/posts/archives/jan2000/949098457.Ch.r.html" target="_blank">http://www.madsci.org/posts/archives/jan2000/949098457.Ch.r.html</a>

Now that's funny. I never would have thought there were two competing theories on this subject. Such a trivial topic as ion structure.

Thanks for that link, it was interesting. I haven't done inorganic chemistry in a long time. I'll have to read some more to freshen up, heh.

P.S. You have replies to some of your threads at your place. :)

PrimoPyro

You have gone way off topic here and I still think sulfate ions have nothing to do with peroxide bonds. I know how covalent bonds work btw, thank you. Please don't treat me like some fucking kewl.

And I still think structure A can't exist under any condition and this can't be explained by sulfate bonds.... :rolleyes:
Please come up with some proof and not something trivial which has nothing to do with it.

<small>[ August 09, 2002, 12:47 PM: Message edited by: vulture ]</small>