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Corvettaholic

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hydrogen peroxide & light
« on: June 04, 2004, 03:44:20 PM »
On another thread, I read that light will decompose hydrogen peroxide. Why? Does it take that tiny little bit of energy to start some kind of reaction? Also, my naming skills are kinda rough, so is hydrogen peroxide HO, OH, or is it something else?

Offline jdurg

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Re:hydrogen peroxide & light
« Reply #1 on: June 04, 2004, 04:26:56 PM »
Peroxide describes an oxygen-oxygen single bond, so any peroxide compound has a something-oxygen-oxygen-something bond.  In the case of hydrogen peroxide, it's H-O-O-H.  Peroxides are very unstable, so a tiny bit of UV light will cleave the O-O bond forming two radicals.  These two radicals will then cleave other peroxides thus causing the entire thing to "go bad."  (Over the long run, it winds up decomposing into oxygen gas and water).  
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Corvettaholic

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Re:hydrogen peroxide & light
« Reply #2 on: June 04, 2004, 04:30:42 PM »
I'm visualizing an effect a lot like fission. I know O-H is unstable, cause I remember drawing octet diagrams of that stuff, and has a hole that needs to be filled. Can you make peroxides with other stuff besides hydrogen? Easily? So, lets say you have a gatorade bottle full of very pure hydrogen peroxide, and you shove it into a tanning booth. As soon as the first O-O bond breaks, the resulting O-H will shoot off and take out two more hydrogen peroxide molecules, and exponentially increase. Man, that would really ruin your day!

Offline Mitch

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Re:hydrogen peroxide & light
« Reply #3 on: June 04, 2004, 04:33:37 PM »
Peroxides are notoriously explosive.
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Offline jdurg

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Re:hydrogen peroxide & light
« Reply #4 on: June 04, 2004, 04:36:26 PM »
Peroxides are the organic chemists, and the alkali-metal chemists, arch-enemies.  They form VERY easily on a lot of oxygenated organic molecules, and on strong reducing agents.  They can ruin your day when you have a mixture with peroxides in it and you accidentally let it dry.  Ethers are the most infamous of all organic compounds when it comes to forming peroxides.  As hmx can back me up on this, a large number of peroxides are classified as very dangerous high explosives.
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Corvettaholic

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Re:hydrogen peroxide & light
« Reply #5 on: June 04, 2004, 05:01:56 PM »
How does drying a solution result in peroxides? I figure they would be there regardless of the state of the compound. Or maybe they're just unstable when not in an aqueous solution? Is there a way to predict peroxides when mixing solutions together? The last thing I want to happen is accidentally blow myself up (which is why I run my stupid ideas past you guys first).

Offline jdurg

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Re:hydrogen peroxide & light
« Reply #6 on: June 04, 2004, 10:24:34 PM »
Organic peroxides are VERY shock sensitive, so when they are in an aqueous solution they are less "dangerous" since they can't really be shocked or experience friction.  When you take an aqueous solution which has peroxides in it and dry it out, the water buffer it once had suddenly dissapears and it begins to crystalize out of solution.  Once it has formed into a crystal, it can then be shocked, rubbed against, moved, etc. etc. which will set off its decomposition.  Peroxides won't form unless the conditions are right for their formation.  For some substances, like diethyl ether, it doesn't take much for peroxides to form.  These substances are generally not available to joe-shmoe, and wherever you could get them from they would be required to provide extensive information about the possible formation of peroxides.  
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Offline hmx9123

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Re:hydrogen peroxide & light
« Reply #7 on: June 05, 2004, 06:20:59 AM »
First, hydrogen peroxide is: H-O-O-H.  Remember that list of bond dissociation energies that I posted?  Take a look at the H-O bond strength and then the O-O bond strength.  Which one would you expect to be cleaved first?  Yeah, the O-O single bond.  I think it's like 35 kcal/mol or something like that.  It's quite weak.

Second, your tanning booth experiment: pure H2O2 would most likely explode.  It's not that stable in highly concentrated form as it is, slowly decomposing at room temperature (mostly due to impurities), but in a tanning booth, it'd probably explode.  If not, you'd at least have a lot of effervescent decomposition.

Next, organic peroxides, in almost every case (I say almost every because although I can't think of an example where this is not a case, I'm sure that AWK will come along and prove me wrong) are dangerous explosives.  Unless you are using hydrogen peroxide, alkali peroxides, ethers, or some fancy organic molecules, you're not that likely to form them.  It's a problem in organic chemistry because we have to evaporate things to dryness and if you've got peroxides in the solution, they don't like the heat of a rotovap bath. :)  Even in solution, peroxides can be unstable, but are generally more stable when in solution.

The military generally won't touch peroxides because they are way too sensitive.  Tiny amounts of friction, flame, shock, heat, or impact can cause them to detonate.  Terrorists like to use them because they are easy to form if you have the right chemicals, but those idiots don't care if they lose a few people making destructive devices.  Hell, they tend to blow themselves up most of the time anyway.  We used to make very tiny quantities of organic peroxides for class demonstrations back where I did undergrad, and they were pretty impressive.  We'd have only about 0.07g of peroxide, less than the size of a crushed up aspirin, and we'd hit it with a hammer, pretty much just letting the hammer drop on the pile, and it would detonate.  The students loved it, but we were careful to be safe about it all.  The remainder of what was left on the filter paper we'd burn off, as it wasn't enough to detonate, and we'd get a resultant fireball.  My avatar is a picture of one of our burn-offs of a nearly microscopic amount of peroxide.  We weighed that particular sample, and it's 0.001g.  That's 1 mg.  Nothing.  You could barely see any white dusting at all on the filter paper, and when we brought a glowing wooden splint up to it, we got a fireball about 10" in diameter.  That is dangerous, I don't care who you are.  So, bottom line: if you make organic peroxides, in whatever capacity, you're asking for a whole lot more trouble than you bargained for, even if you think that you're prepared for anything.

Corvettaholic

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Re:hydrogen peroxide & light
« Reply #8 on: June 05, 2004, 05:42:17 PM »
Wow, peroxides sound like a huge pile of no fun. At least if fun involves staying alive anyway. I think I remember one of jdurg's alkali articles going over peroxide formation when dunking alkali metals in water? How does that work? Take potassium for instance, K + H2O = some peroxides? Well since they DO form, I can imagine why alkali metals lower on the periodic table are very scary.

Offline hmx9123

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Re:hydrogen peroxide & light
« Reply #9 on: June 05, 2004, 08:56:24 PM »
Actually, potassium peroxide and superoxide, for example, are inorganic peroxides and much less dangerous than their organic cousins.  They are still powerful oxidizing agents to be sure, but they by themselves are generally not horrifically unstable like organic peroxides.  And, yes, the alkali peroxides just form by oxygen reacting with alkali metals, although most of it reacts to form the oxide initially.  You kind of have to control your conditions to get the peroxide or superoxide to form.

Corvettaholic

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Re:hydrogen peroxide & light
« Reply #10 on: June 05, 2004, 11:08:08 PM »
Ok, I think I'm a bit fuzzy on the types of peroxides. As I understand it, a peroxide is H-O-O-H, plain and simple. So whats the difference between organic, inorganic, alkali, and other types? I'm guessing it involves replacing the hydrogen atoms with something else, so long as the link between is ?-O-O-?. Am I on the right track?

Offline jdurg

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Re:hydrogen peroxide & light
« Reply #11 on: June 05, 2004, 11:28:45 PM »
Yes, you are on the right track.  A peroxide is classified as a compound with the structure X-O-O-X where X can be any element or compound.  An organic peroxide is just a peroxide where the Xs are organic groups.  Alkali metal peroxides have an alkali metal in place of the X, and hydrogen peroxide has hydrogen atoms there.  The heavier alkali metals are indeed dangerous because their larger atomic radius makes it much easier for them to form peroxides and superoxides.  (The larger radius gives them a bigger electron cloud which allows oxygen to bind much easier to it since it's not repulsed by the nucleus of said atom).

The peroxide doesn't form on contact with water.  What happens is the peroxide forms on the surface of the metal as it reacts with atmospheric oxygen.  The further down the column you go, the less oxygen is required for the peroxide formation to occur.  (Lithium won't form peroxides, sodium will with some effort, potassium will with minimal effort, and rubidium and cesium form them without any effort at all).  The danger with the alkali metal peroxides is their decomposition in water.  The alkali metals release a lot of energy when they react with water, and any peroxides on their surface only add to this release of energy.  So when the alkali metal reacts, it causes the peroxides to decompose as well, and these generally do so with great force and suprise people.  Throw a small piece of potassium in water and you can have an idea of what to expect in terms of a reaction.  Now take that same sized piece but have it covered in peroxides and the explosion will be much more intense and will usually surprise the person who carried out the reaction.  Explosions are dangerous, but even more so when they occur unexpectedly.
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