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Author: Subject: atmospheric combustion of Al (foil)
Polverone
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[*] posted on 25-9-2004 at 18:52
atmospheric combustion of Al (foil)


When I was much younger I used to make a variety of simple pyrotechnic devices with KNO3 as oxidizer and whatever fuels I could access. Sometimes I used aluminum foil as casings for my mixtures, and I remember being pleased by the intermittent ignition of aluminum that would briefly produce very bright white flames/sparks (but the combustion never lasted long).

Recently I bought some very cheap aluminum foil for kitchen use, but it turned out to be very thin and the box had a poor cutter design so it was relegated to the lab. I then decided to try to find a way to reliably ignite this aluminum foil. The oxidation of Al to Al2O3 in atmosphere is very vigorous and showy but difficult to initiate due to the tenacious oxide layer that forms on all aluminum.

I first tried wrapping aluminum foil around large, lumpy crystals of NaNO3 so that there was a lump of nitrate with extra foil around it at one end, with a sort of loose tube of wrapped foil extending up to where I held it with pliers. Heating such devices in the flame of a propane torch, I was able (with difficulty) to obtain ignition maybe 1 out of 2 trials. Curiously, I tried and failed with KNO3 several times. I don't know if it had something to do with chemistry or more with the fact that I started with large single broken crystals of NaNO3 and only small prills or powder of KNO3.

I tried a number of other oxidizers to obtain easier foil ignition: KMnO4, KClO3, and Ca(OCl)2. I was not successful with any of them. I also tried using zinc powder and a mixture of NaNO3 and sulfur. None of these worked very well. Adding small quantities of NaOH or Na2CO3 to the NaNO3 didn't seem to help (I thought alkali might help to expose bare metal).

Finally, I found one improvement that seemed to lead to more reliable ignition: adding a matchhead-size lump of charcoal to the package wrapped up in the tip with the NaNO3 almost always leads to foil ignition when the assembly is heated in a torch flame.

Sometimes the flame burns out halfway up the foil; other times it consumes the foil right up to the pliers. It seems to burn best when the foil is rolled up into a loose tube. The burn is dazzling and fast, over in less than four seconds with my ~10 cm lengths of rolled up foil. If I were a chemistry teacher this is one demo I'd love to show my students, since seeing a household item like aluminum foil undergo such a vigorous reaction is arguably more impressive than doing the same with magnesium ribbon or powdered aluminum from the stockroom.

I'd love to find a way to guarantee that the entire mass of the foil would be consumed once ignition is achieved. I may experiment with different geometries for the foil. I don't know if thicker foil would be better or worse for easy ignition and sustained burning. I will also have to try KNO3 again with the aid of charcoal.




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[*] posted on 25-9-2004 at 20:45


Have you tried igniting the foil with the help of magnesium ribon? I'd bet that the best geometry would either be with very thick foil or very thin foil. Thick foil would dissipate heat too quickly, although if you get very thick foil burning, the amount of heat being generated might make up for this. This would be like those aluminum warships that started burning and couldn't be put out.
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[*] posted on 25-9-2004 at 21:29


No, I haven't tried Mg ribbon. That might be a yet more reliable way of doing this. I tried a few burns with heavier aluminum foil, and it seems tougher to ignite but also burns better once started. I wonder if an oxy-acetylene torch with an oxidizing flame would do the job.



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[*] posted on 26-9-2004 at 04:06


Would it be possible to prime the surface of the foil first. Maybe wipe it over with sodium hydroxide and then with nitric acid to give a layer of nitrate and oxide rather just oxide.

I though I heard during the Falkland war that it is a bit of a myth about structural aluminium catching fire because it softens and melts so quickly. It only readily burns when it has a large surface area. I though the fires were realy when it gives spectacular flashes when ordinancy detonates (you have already got very serious problems by this time).
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[*] posted on 26-9-2004 at 07:03
Al Foil


The problem does seem to be that pesky oxide layer. I tried it by placing a
thin coat of superglue on the foil and immediatedly pressing it into
some finely powdered KClO4. After that I rolled the foil up and ignited with a
propane torch. It took several attempts to burn this roll down completely.




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[*] posted on 26-9-2004 at 08:05


Try this:

Leave your Al foil in an alkalicarbonate solution for a few minutes, you'll notice the H2 production and what's more important, the Al isn't shiny anymore. Now if you quickly light this after taking it out of the solution you should have more succes.

Why not with sodiumhydroxide? Because the reaction with carbonate takes days to dissolve all the Al.




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[*] posted on 26-9-2004 at 08:15


Old flash bulbs contained aluminum wire, sometimes with an ignitor or some other material, surrounded by pure oxygen. Some bulbs do contain magnesium or zirconium wire, I don't know which contained what. It is a fact that one bulb flashing will set off another bulb right next to it from the UV light! I have actually tried this a few times, but not recently, as old flash bulbs are getting scarce. I have also made Al foil flash in oxygen by passing current through a strip of foil while in a container. The O2 was generated by H2O2 catalyzed by permanganate. It was a long time ago, and I'm not sure if the foil was completely consumed, or the exact concentration of the O2. I may have to repeat the experiment, varying the pressure of the O2 and the thickness of the foil. Maybe it would even burn in regular air if the pressure was higher? BTW it is quite easy to get Al to slowly oxidize in air; just put a little Mercury on it, and rub. It will get very hot, making white Al2O3 and lots of Hg vapor. I found this out in the 5th grade while carrying some Mercury in a gum wrapper, in my pocket. I'm sure all my old schools are toxic waste sites now, as all the kids used to put mercury on pennies, silver dimes, and rings, and of course, the floor.
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[*] posted on 26-9-2004 at 08:27


Try using thin foil in the beginning to ignite easier, then have that continue into thick foil for better burning (glue two sheets of thin and thick foil edge to edge for this).



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[*] posted on 26-9-2004 at 12:03


If you put an aluminium block to soak in stuff called decon 90, I think it is a tetra-alkyl ammonium surfactant, it will very slowly bubbles after a while and the aluminium ends up nearly black, which is hard to remove. I have always wondered what it is or whether it is more passivating than the oxide layer. It might be due to stuff in the aluminium to make it harder because aluminium does not usually give black compounds.
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[*] posted on 26-9-2004 at 12:29


Thanks for the suggestions. I had thought of trying to ignite the foil in pure O2 instead of air, but that involves somewhat more work. I'm sure it's spectacular when executed correctly, though!

I did consider treating the foil with something that strips the oxide layer but I don't think it will work very well. I need to dry the foil before igniting it, and by that time I think it will have reformed the oxide layer. Still, it might be worth a shot.

It seems important that the internal combustion suddenly expose fresh, hot aluminum to the air for ignition to succeed. This means that the glowing blob at the tip, laden with molten nitrate, aluminum, etc. falls off (often commencing spectacular combustion of its own) and the remaining mass of foil then begins burning. The way burning peters out (when it fails to consume most of the foil) usually seems to be that the foil melts down to a blob at the tip and the reaction stops self-sustaining. Once the reaction slows, even quickly exposing it to the torch flame again cannot make it resume. I don't know if there's any reasonable way to make it harder for the foil to melt down to an unreactive glob at its tip.

Also spectacular in a slightly different way is strongly heating Al foil with NaOH in atmosphere until brilliant sodium-aluminum fire begins. I was inspired to try this by the many posts on sodium "thermite" mixtures, and it is actually what made me think of trying to ignite aluminum foil by itself. I should try using NaOH as the fire starter at the tip of some foil, but handling NaNO3 is much easier.




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[*] posted on 26-9-2004 at 14:17


If you could find a flux that would make the aluminium wet the other stuff, a bit like soldering.
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[*] posted on 26-9-2004 at 15:17
Al and Hydrogen


It's funny that my Senior year chem teacher didn't believe me when I
said hydrogen could be produced from the reaction of NaOH and Al(1975).
But what the hell did I know, I don't have a PhD in chemistry - except that
I had been doing it for 4 years before I took her class. Good ole Drano and
aluminum foil gave me the H2 I needed for my flammable balloon
experiments ! I did experience an awesome flash when I stole the KClO3
and red phosphorus from her lab ! Lost my eyebrows BTW !
When it comes to lab safety - I THINK SHE HAD A VALID POINT !

I've been focused on both foil and powdered Al since. If you can get it to
burn in open air with a little help I'd like to know about it ! Powered Al -
no problem, foil Al definite problem !

[Edited on 26-9-2004 by MadHatter]




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[*] posted on 26-9-2004 at 15:48


Of course it can. Your chemistry teacher apparently did not know that Al, or Al2O3, is amphoteric, i.e. forming both cations (hydrated or partly hydrolysed Al+++) and anions (AlO2- or AlO3---, or Al(OH)6--- if hydrated) with strong mineral acids and alkalis respectively. H2 being evolved in both cases when the metal is used.

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[*] posted on 27-9-2004 at 03:29


Burning Al is fun.
I have not tried burning the foil but I have successfully burned Al cans before.

Take your Al can (it will be labelled 'Aluminium - Recycle' somewhere on it), clean out any crap and then cut off the top third. Fill about halfway with your oxidiser of choice. I use Sodium Nitrate for several reasons:

melts easy
not toxic
not that vigourous a reactant
cheap and easy to get

Then add something combustible. I used twigs and small sticks stuffed into the nitrate but anything that won't react all in one go is fine. Give that mixture a good stir then place on hot coals / ashes. Or place on ground and light fire around it.

The nitrate melts and reacts with the combustibles present. This usually gives out enough heat to start the reaction of the nitrate with the Al can itself. Once it starts there is no stopping it. It will melt the rest of the can in an amazingly bright (don't look directly at the white fire) deflagration. Once it has started you can, however, chuck more Al cans on and maybe, if you're lucky and haven't been blinded, end up with an Al fire burning in atmospheric Oxygen only. Once the can goes up you could probably chuck Al foil on too.

If you get it right is is spectacular - if you don't then all you get are a few white sparks and a sense of disappointment.
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[*] posted on 27-9-2004 at 09:19


You could also shred the top third of your can (or even some more cans) and add to the nitrate-reducing agent mixture.
I suggest this: a mixture of coarse charcoal powder, aluminium bits and NaNO3 in an Al can. Use a lot of Al bits, for you also have at your disposal atmospheric O2 (I think this method will work for sure to ignite Al in atmospheric O2).




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[*] posted on 27-9-2004 at 10:03


You guys are using so much oxidizer/pyrotechnic mixture that it's hard to judge if it's indeed a burn in atmospheric oxygen...<_<



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[*] posted on 27-9-2004 at 16:55


What doesn't work. I tried to get Al foil to burn in pure O2 at atmospheric pressure and all I got was a short flash. I heated a 1 cm by 15 cm strip of foil with a 12 volt car battery. The foil was fairly thick and carried a lot of current before it melted in the thinnest part. I then tried to ignite it in the pure O2 (from a welding tank), using a 12KV neon sign transformer. The spark jumped the gap, but only melted the ends of the foil with hardly any brightness. Things to try yet: Thinner foil and higher than atmospheric pressure O2.

[Edited on 28-9-2004 by Mr. Wizard]
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[*] posted on 28-9-2004 at 09:55


Isn't commercially available Al foil coated with PE for strength? The decomposition of this is rather endothermical...



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[*] posted on 28-9-2004 at 11:03


One side of the foil is rather dull, and may have a PE coating. It was also hard to get a good connection with the alligator clips. A little warming with a flame, and a sniff test will tell me if if it has PE or some plastic on it. I don't think it does, as the use of foil in an oven would lead to a house full of smoke. There are other plastics that will hold up in an oven though, as the turkey roasting bags show. More tests ahead.
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[*] posted on 28-9-2004 at 19:46


To get rid of the oxide coating and make the aluminum more reactive you could try a small amount of mercury. Wash initally with an organic solvent to remove the polymer layre on the other side and you might have some highly reactive aluminum to show for it.



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[*] posted on 28-9-2004 at 19:56


Even if you could get the aluminium foil to burn after your proposed mercury treatment, you would risk poisoning yourself with mercury vapor. Besides, one would think that that the Hg would readily form an amalgam with the Al (most similar low-melting metals do, e.g. Zn, and also Ag and Au, the basis of an old method of Au refining), while not affecting Al2O3 at all.

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[*] posted on 28-9-2004 at 21:14


I didn't mean to actually "burn" the Al foil with Mercury, it just forms an amalgam as you stated. The amalgam, which doesn't have the benefit of a protective oxide layer, reacts spontaneously with the O2 in the air and forms Al2O3 as a white powder, and lots of heat, and no doubt lots of toxic Hg fumes. If the Hg doesn't oxidize or evaporate, it will continue to catalyze the oxidation of the Al.

Sorry about the confusion, I was just mentioning the Al on the gum wrapper to illustrate how active Al is, not to recommend Hg use.
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[*] posted on 29-9-2004 at 20:04


Your mention that NaNO3 worked and KNO3 didnt makes me wonder if the water from the NaNO3 had an affect since it would likely contain much more water than KNO3. Really neat discovery by the way.
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[*] posted on 29-9-2004 at 23:02


As far as I can tell, the NaNO3 is anhydrous. It was evaporated to a dry solid under heating (made in solution from NaOH and NH4NO3) in a pyrex dish and chipped out, crushed, and bottled while still warm. There is a small amount of excess NH4NO3 present, detectable by odor of ammonia with NaOH solution; I don't know if it's enough to affect things. Unfortunately, I do not have any purchased NaNO3 to make a comparison with. I'd have to grow crystals to get purer material.

I gave it four more tries tonight with KNO3, 3 times with charcoal and once without. All were failures. I tried it twice with KNO3 with ~20% NH4NO3 mechanically incorporated plus charcoal, and those too were failures.

According to my Lange's Handbook of Chemistry, NaNO3 density is 2.257, mp 308 C, decomposes at ~380 C, while KNO3 density is 2.109, mp 334.3 C, decomposes at ~400 C. If we take the amount of oxygen per unit volume to be one for KNO3, the "oxygen density" of NaNO3 is 1.27. In addition, the lower melting and decomposition points of NaNO3 might help explain why it starts the aluminum combustion better than KNO3. I can try to rationalize the behavior I see after the fact, using known properties of the substances involved, but I'd have to say this difference in foil-lighting-capability is pretty unexpected. I will have to give lithium nitrate a try one of these days (but I'll have to work quickly to keep the LiNO3 dry).

I've done my ignition trials over an open basin of water (though dry sand might be better) and I've noticed something else delightful and unexpected: at least some flaming pieces/drops of aluminum that fall from the foil will "skate" around just above the surface of the water until they burn out. This can last up to 2-3 seconds.

This sort of investigation is thrilling yet inexpensive, unmotivated by theoretical or commercial needs, and basically detached from the world of Official Chemistry. It reminds me of some of my most-enjoyed times as a child, learning about the contents of household products or the behavior of pyrotechnic mixtures through hands on investigation. I wonder if I could get this published in the Journal of Chemical Education with enough refinement of procedure and added theory. "A demonstration of the rapid atmospheric oxidation of aluminum suitable for student reproduction with household materials." :P

[Edited on 9-30-2004 by Polverone]




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[*] posted on 3-10-2004 at 04:08


I've heard that liquid gallium affects aluminum in the same way
that mercury does. Has anyone ever tried this? Although Ga is
much more expensive than Hg, Ga is virtually harmless.
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