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Author: Subject: Exotic thermites & analogs
Theoretic
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shocked.gif posted on 16-12-2003 at 01:34


Just an interesting fact: Fe/Fe2O3 releases more energy per gram (upon conversion to FeO) than CuO/Zn!
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[*] posted on 16-12-2003 at 02:42


maybe so, but that's definitely not the case with Al!! :D:D



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[*] posted on 16-12-2003 at 13:05


Can anyone explain the reaction between Al and S? There is not any oxygen present so I would think the heat oxidises one of those and then they react. Not sure thought.



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[*] posted on 16-12-2003 at 13:15


S can act as an oxidizer. An oxidizer is a substance that can increase the oxidation number of other compounds, by decreasing it's own.

In this case, the S will be reduced to S 2- and the Al oxidized to Al 3+.




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[*] posted on 16-12-2003 at 20:32
CuO homemade vs 'store bought'


Maybe Chemoleo's 'drying' of his home made CuO at a red heat, had converted some of it to the Cuprous Oxide? If he heated it in an oxygen atmosphere wouldn't it stay at the higher oxidized state?

[Edited on 17-12-2003 by Mr. Wizard]

[Edited on 17-12-2003 by Mr. Wizard]
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shocked.gif posted on 17-12-2003 at 01:37


Chemoleo:
Hell, no! :o:o:o
Fe3+ => Fe2+ has a higher potential than Cu2+ => Cu! Not Fe3+ => Fe though, :D:D:D.
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mad.gif posted on 17-12-2003 at 07:05


Theoretic:
what are you getting at? Per gram of Al, the oxidation to Al2O3 is surely releasing more energy than with Fe or Cu, as the main factor limiting this is the atomic mass (moles per gram) of the metal. That's all I suggested, so pls don't write posts on different issues putting it as if I was talking bollocks :o:o

Mr Wizard - I doubt it. Cu2O should oxidise to CuO in an oxygen atmosphere, under such conditions. Anyway, dissolving the homemade CuO in HCl produced green CuCl2 only, and no white precipitate which would indicate the presence of Cu(I)Cl!
I think I will one day try to obtain CuO by different means, i.e. not by the hydroxide/heat method

[Edited on 17-12-2003 by chemoleo]




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sad.gif posted on 17-12-2003 at 07:48


Chemoleo:
Sorry, I thought that you were talking about Al2O3 as a fuel in Fe2O3 and CuO thermites.
You want to use methods other than the hydroxide method for CuO? Can't think of any ones apart from heating copper nitrate and copper sulfate, both of which produce nasty fumes.
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[*] posted on 19-12-2003 at 21:21


yeah nasty but I wouldn't use such disapproving word for (holy) SO3:o
got some sodium dichromate lately. may try CrO3-Zn. CrO3 is a little scary but not that much (?) I think it's OK as far as I don't let it touch org matters




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[*] posted on 20-12-2003 at 10:21


haha... I was thinking on the lines of CrO5 (a beautiful blue compound), however! I bet that explodes/reacts very violently upon touching a metal such as Al!



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[*] posted on 20-12-2003 at 12:07


IIRC copper carbonate will decompose to CuO and carbon dioxide with heating. Copper carbonate is also available from many pottery suppliers BTW. If you bother going to a ceramics store then I suppose you might as well get CuO though.

Chemeoleo, you describe your home made CuO as quite bulky and free flowing, right (flows almost like a liquid)? I think this may be why it doesn't work as well as the more dense, store bought stuff. I think all thermite like reactions, including CuO/Al, are solid state, compared to the gaseous state of most pyrotechnic compositions. This means that there is no gasseous flame front to travel between the particles, but rather that the particles must pretty much be touching to react. The result is that if you bulk up a thermite composition, it will burn slower. If you pack it down it should burn faster.

To test this you could take some of your CuO/Al mix with home made oxide. Pack it carefully into a tube, then compare it to mix using commercial stuff, packed to the same density.
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[*] posted on 20-12-2003 at 13:41
WoW!


This is great stuff!!!
It gives gas volume, cal/g or cm3 etc etc, temperatures for every thermite reaction imaginable, plus many more reactions!!

Check where Al/MnO2/CuO/PbO2 sits!

Do you have some more data files like this?

Edit: Sparky, I agree it's a parameter, but I try to maximise density for any pyro mix.. i dont think that was the reason anyhow. Yes, I think I will make some CuCO3 (this is what I had in mind originally, referring to Theoretic's post), and heat it until it turns thoroughly black. Will test this again. It did baffle me as you can see! :(

[Edited on 21-12-2003 by chemoleo]




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sad.gif posted on 3-1-2004 at 19:31


the file HAS BECOME corrupted. something's going wrong with the attached zip files (?) ...



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[*] posted on 26-1-2004 at 16:48
Mad Science Thermites :D


Well, I did some interesting, but slightly mad stuff yesterday.
It turned out I had quite a bit of silicon (the paste type for sealing things) lying about, so the decision was made to burn it. It burns not great, but once it does, it develops lots of white smoke and leaves a grey white substance - whcih is carbon (& derivatives)-contaminated SiO2 I speculated.
So I went about testing this :) . The residue from the burnt silicone was crushed, and mixed with an estimated, not weighed amount of Al 200 mesh. It was started with a film capsule full of NaClO3/Al. After the initial flash of light from the NaClO3/Al, the silicone burn product/Al indeed started to glow a bright white glare, the reaction took several minutes to subside! Chunks formed, which glowed bright white and could be separated from the rest!
Today I looked at one of those chunks, they turned black/brown that crumble easily... Nonetheless purifcation of pure silicone shouldnt be easy, my textbook states that boiling in conc. HF is necessary :(
Nonetheless, cool experiment isnt it? ;)

Actually some liquid silicone was mixed with Al powder, and then allowed to harden. I was hoping of course that once I wuld light it, the resultign SiO2 would react with the Al, make a thermite paste of sorts :) - but sadly, although this burned better, no thermite reaction occurred.

The second one involved ammoniumdichromate. This was alighted, and the resulting extremely voluminous Cr2O3 was used for a thermite reaction, with Al again. It works very nicely, and produces a solid product where no free Cr could be found. In fact it completely maintained the shape of the plastic cup in which it was ignited (even though it was glowing white for at least a minute)... was thinking that this could have all sorts of useful purposes!!

Anyway, i think it's time to purchase some fine quartz sand from pet shops etc, and see how the real SiO2 works!

Edit: Post No 300 in the true mad science way :D:D

PS what happened to Adiabatic? The guy who posted this thermite reactions file? His post completely disappeared!?!

[Edited on 27-1-2004 by chemoleo]




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[*] posted on 26-1-2004 at 17:28


Quote:
Originally posted by chemoleo
The second one involved ammoniumdichromate. This was alighted, and the resulting extremely voluminous Cr2O3 was used for a thermite reaction, with Al again. It works very nicely, and produces a solid product where no free Cr could be found.

I don't quite get it. Are you saying the product was something besides Cr itself?

Quote:

Anyway, i think it's time to purchase some fine quartz sand from pet shops etc, and see how the real SiO2 works!

If you have water glass (sodium silicate solution), you can also precipitate fine SiO2 by adding some HCl or other acid, then rinse and dry. Kind of an expensive way of making sand, but it is pure and fine that way.

I've tried a couple of thermites myself recently. I've made iron oxide in a couple different ways (using FeCl2 and FeCl3) and mixed with aluminum in the appropriate proportions. Actually, the mixture made with iron (2) oxide seemed to react the fastest.

I made some CuO and made some thermite using that. Although the reaction was pretty fast, I didn't get a flash or a large plume of Cu vapor (I only used 1g total of the thermite). I did find a small ball of solid Cu. Maybe my CuO wasn't completely dry, although I did heat it at 200C for about an hour after all the water was gone.

I made some thermite using ZnO and Al as well. This is kind of weird stuff. It sort of glows for a while, then suddently flashed up with a green flame, then glows for a while and flashes again. Very little Zn was produced - I'm thinking it either vaporized or else burned in the air as soon as it was formed. There was however a thin coat of Zn on the metal ignited it on. Probably condensed vapor.

I tired a mixture of CaSO4 and Al. With hydrated CaSO4, the reaction is fairly fast. When I dry the CaSO4 first, it is like a flash powder. I expected to get maybe CaS but treating the residue with water produced no gas. My guess is it went all the way to CaO.

I've tried a couple of different ways of igniting the thermite mixtures. Keep in mind that I am doing this in my sink inside so I am using small amounts (about 1g) and I don't want something that produces a lot of SO2 or other irritating gases. The best thing I have found for igniting is a small amount of fine Al powder and S. This ignites reliably with a "fuse" made of string soaked in KNO3, and it always sets off the thermite (even the CaSO4 mixture). Very simple, and doesn't produce much in the way of fumes as long as I don't let the Al2S3 sit too long before disposing of it down the drain. I tried a sugar/KNO3 mix and it lit with the string fuse but would not reliably ignite the thermite (even iron/aluminum). I tried a thin nichrome wire heated with a battery, and this was not reliable either unless I used it to first ignite the Al/S mix. Magnesium ribbon works most of the time but is hard to light even with a butane torch and I'm afraid I will light the thermite in the process of trying to get the Mg to light so I don't usually use this method.

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[*] posted on 27-1-2004 at 15:34
Zn thermite


I recently came into contact with 11lbs of 625 mesh Zinc powder and i have some MnO2 and can make some Red Iron Oxide....I was just wondering if any of u knew and stoichiometric ratios for either of those mixes. any help is greatly appreciated



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[*] posted on 27-1-2004 at 18:40


Quote:
Originally posted by FrKoNaLeaSh101
I recently came into contact with 11lbs of 625 mesh Zinc powder and i have some MnO2 and can make some Red Iron Oxide....I was just wondering if any of u knew and stoichiometric ratios for either of those mixes. any help is greatly appreciated


Using the zinc in place of the aluminum, I assume? Let's see.

Fe2O3 + 3Zn --> 2Fe + 3ZnO
160 195
That would be 160 parts Fe2O3 to 195 parts Zn, or roughly a 4 to 5 ratio of Fe2O3 to Zn.


MnO2 + 2Zn --> Mn + 2 ZnO
87 130
87 parts of MnO2 to 130 parts Zn, or roughly 3 to 4 ratio of MnO2 to Zn.

I did some energy calculations - may or may not be right so I won't show them here. But if they are right, I show only 150 calories per gram for the Fe2O3 mixture and 195 calories per gram for the MnO2. By contrast, the standard thermite Fe2O3/Al reaction yields 930 calories per gram. My ZnO/Al mixture burned rather poorly and it yields 488 calories per gram. So unless my calculations are wrong, my guess is using zinc in place of aluminum is not going to work very well.

But its worth a try to see what happens. Always use small amounts and stay well back until you have seen what a new reaction does.

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[*] posted on 27-1-2004 at 20:25
thank you


omg u are the best thank you so much.....i am not the best at stoichiometry yet i am not learnign it till next week in chem class.......and no one on totse knows anything like that so thanks

and also i am not exactly looking for an insane reaction....i have never seen any thermite to compare it to except a pic in my textbook. so i will try both and report back what i find....also i do have 11 pounds of Zn to use...so i might as well experiment

[Edited on 28-1-2004 by FrKoNaLeaSh101]




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[*] posted on 29-1-2004 at 20:03
I don't think this worked


Tried making thermite with SiO2 and Al. I have heard that this works. I took 0.5g of fine play sand ("asbestos-free";) and 0.3g of pyro grade Al powder. Attempted to ignite with a mixture of 0.3g S and 0.3g of the Al powder (this ignition always works for other thermites, and even works when using CaSO4/Al). It did not ignite. BTW, I know the proportions are not right for the Al/S. I use excess Al to prevent SO2 from being formed since I'm doing this inside.

So I tried mixing some of the aluminum/sulfur mix in with another batch of the SiO2/Al thermite. This time the mixture did burn, though rather slowly.

When cool I added a few drops of 3M H2SO4. I expected the production of silane, which burns spontaneously in air, from the aluminum silicide that is formed as one of the products of the reaction. But I got no flames or pops - only some H2S. Then I added a bit of water and let the Al2S3 react. Once the reaction stopped, I added more dilute H2SO4. Again, just a few more bubbles of H2S. The resulting solution was black, not brown like I would expect for silicon.

Has anyone else here gotten this to work? I've seen several posts on Usenet where people claim to produce Si using this reaction.

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[*] posted on 30-1-2004 at 13:41


quote:
I expected to get maybe CaS but treating the residue with water produced no gas. My guess is it went all the way to CaO.

the realy realy hot CaS dispersed in air couldn't survive and probably "burnt" to produce chalk....

btw anhydrous CuSO4-Al and PbSO4-Al should make pretty violent thermites. not something to do in your kitchen specially the last one :P
3CuSO4 + 10Al    > 3Cu + Al2S3 + 4Al2O3




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[*] posted on 7-2-2004 at 16:08
Success! (SiO2 Thermite)


I found out that the problem with the SiO2/Mg thermite was that the sand I was using had too large of a grain size. I was able to find some 200 mesh "silica", and was successful with this.

I mixed 1.2 grams of SiO2 and 1 gram of powdered Mg. I ignited this with a small amount of Al/S mix. It glowed red hot for almost a minute. When it cooled, I dropped a small amount in dilute HCl. I thought I heard a couple of pops, then I definitely saw a small spark on top of the solution and heard a small pop. I then dumped in the entire product into the HCl. For a second or two, it sounded like firecrackers going off (though not as loud). There were numerous flashes of flame. After that there continued to be pops and flashes every few seconds for a couple minutes.

I washed the resulting product and crushed it up. Presumably it is impure Si. It has a color between gray and tan and looks a lot like mud. When viewed from the outside of the glass measuring cup a metalic shine is aparent where the surface is in contact with the glass.

I took some videos and may post still shots of the reaction and products sometime depending how the still frames look once I digitize them.

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[*] posted on 11-3-2004 at 19:14
magnisium oxide?


Would it be possible to create a thermite from magnesium oxide and aluminum yielding magnesium?
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[*] posted on 12-3-2004 at 08:30


Yes, but you would get vaporized magnesium in the reaction flame.....

Check the boiling point of Mg...




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[*] posted on 12-3-2004 at 09:55


Quote:
Originally posted by vulture
Yes, but you would get vaporized magnesium in the reaction flame.....


Are you sure? The reduction potential of Mg2+|Mg is -2.37V while the reduction potential of Al3+|Al is -1.66V Hence, aluminium will not be able to reduce magnesium oxide. Oxides of metals only below aluminium in the electropositivity series will work in thermite mixtures. And the melting point of Mg is 650oC.

[Edited on 12-3-2004 by t_Pyro]
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[*] posted on 12-3-2004 at 15:23


Quote:
Originally posted by t_Pyro
The reduction potential of Mg2+|Mg is -2.37V while the reduction potential of Al3+|Al is -1.66V Hence, aluminium will not be able to reduce magnesium oxide. Oxides of metals only below aluminium in the electropositivity series will work in thermite mixtures.


Actually, it is the heat of formation, and not the reduction potential, that will determine if a thermite mixture will work. Aluminum is able to reduce K2O to potassium even though potassium has a larger reduction potential. But the aluminum loses 3 moles of electrons, whereas the potassium only gains one. Thus that reaction end up being slightly exothermic, and will proceed. OTOH, the reaction between MgO and Al is endothermic if you work out the energies involved. So you are correct that this reaction will not proceed (at least not without adding heat). The reverse reaction (Al2O3 and Mg) theoretically could be used to prepare Al, although the difference in heats of formation is small enough that the reaction might not be self-sustaining.
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