Sciencemadness Discussion Board

Aluminum and sulfuric acid makes...flash??? [pics]

Quince - 1-3-2005 at 06:44

Initially, I was using a blender to make Al powder from foil (in alcohol to minimize oxidation) for thermite. I was unable to ignite the thermite with a sparkler, however (my iron oxide was Fe3O4 from burning steel wool in air). I was messing with the Al, and I used a mortar and pestle on the coarse particles:


I decided to add some H2SO4 (drain cleaner, it's quite concentrated since it starts white fumes right away when heated) to see if H would be produced:


I only had a bit of H2SO4, enough to wet the powder when mixed. It was foaming nicely, and eventually all the powder clumped together instead of falling apart (in later tests I found that more H2SO4 made this not ignite):


I left it while I was doing other shit, but in a couple of minutes I saw some vapor rising from it; it was irritating and like SO3 fumes from heated H2SO4. All the liquid quickly boiled away, leaving a solid foam (note: later I found this reaction could be forced sooner by heating):


It was crumbly, but I managed to flip it over; big bubbles on the underside. It wouldn't light from a propane torch, so I put some pieces from a sparkler on top:


That did it; the light was so bright that you can see in the picture the camera's flash barely illuminates the surroundings in comparison; SO2 smell was produced; the droplets flying around are molten aluminum:


Prying over the aluminum bits showed it had melted into the earthware; later I added water and that caused H2S smell to be produced *chokes*:


Well, I tired figuring this out from the equations, but no cigar. Initially, it's just Al with a bit of Al2O3, so when adding the acid:
2 Al + 3 H2SO4 --> Al2(SO4)3 + 3 H2
Al2O3 + 3 H2SO4 --> Al2(SO4)3 + 3 H2O (I think?)
As I think I do not add enough H2SO4, there is some Al and Al2O3 left.
Then it sits in the air, and after a while something happens (reaction with something in the air?) that heats it up and boils off the remaining liquid, but what? Heating the resulting foamy solid with a torch doesn't do anything, but the sparkler manages to ignite it. Again, what could be going on? Main products of the flash seem to be Al and S2O; a bit of yellowish tint is visible around the aluminum, and when I add water H2S comes off, suggesting perhaps some Al2S3 is also present (Al2S3 + 6 H2O --> 2 Al(OH)3 + 3 H2S). But I really have no idea what could be causing such an energetic reaction. Any ideas from you chemistry gurus?

[Edited on 1-3-2005 by Quince]

Bert - 1-3-2005 at 09:05

Sulfates can be used as oxidizers with Al or Mg to make "flash powder", albeit one with a very high ignition temperature. You obviously made some Aluminum sulfate, perhaps leaving some unreacted Al to be oxidized by the sulfate when you later ignited the mix.

Um, dude, your "bench top" looks quite igniteable...

[Edited on 1-3-2005 by Bert]

guaguanco - 1-3-2005 at 09:07

My head is a bit fuzzy from a cold, but I'd guess that your first reaction was an redox reaction between H+ and Al.
3 H2SO4 + 2 Al -> 3 H2 + Al2(SO4)3

Then later on you managed to ignite the Al metal in air:
2 Al + 3 O2 -> Al2O3

Your pictures strongly suggest this kind of pyrolysis.

[Edited on 1-3-2005 by guaguanco]

chloric1 - 1-3-2005 at 13:12

Um yeh! Pyrolisis of bed covers! You might want to invest in firebricks. They sell six in a box for $7.99 at Menard's where I live.

Quince - 1-3-2005 at 13:58

Initially, the acid turns the oxide coating of the aluminum into sulfate and water, but after the exothermic boiling and drying reaction, there's no unexposed Al metal in the mixture before ignition (it's non-conductive, not to mention that it doesn't look anything like Al).

As I said, most of the product of the burning is Al, not Al2O3. Only a fraction of the Al could have oxidized given the amount of Al left, and I don't see how a bit of the Al burning could have made this much heat. Also, as I said, there were THREE reactions:
1. hydrogen producing during initial acid addition
2. delayed exothermic reaction that boils off remaining liquid, happens only if left exposed to air (no reaction if kept sealed)
3. some sort of burning upon ignition with sparkler

1 is obvious, guaguanco proposed a suggestion for 3, but 2 is unexplained.

Bert, I don't see how the sulfate could have oxidized Al. It sounds like you are proposing something like 8 Al + Al2(SO4)3 --> 4 Al2O3 + Al2S3. Is this likely? If so, that leaves the question of what reaction 2 is, and why I have such an excess of Al at the end (adding more H2SO4 makes the mix unignitable). Also, the smoke you see in the picture smells like SO2. Maybe an additional reaction, Al2S3 + 3 O2 --> 2 Al + 3 SO2? Alternatively, 2 Al + Al2(SO4)3 --> 2 Al2O3 + 3 SO2, but again, the product is mostly Al, not the oxide.

chloric1, this is an old carpet, and the drops of molten aluminum look bigger in the picture than they actually are due to the glow saturating the camera. Most of the stuff didn't fly around as it appears, and instead simply melted into the earthware (when I did a larger amount later, it actually boiled some of the earthware :)).

Maybe someone can repeat this experiment and get a better idea than my descriptions. It works with even coarse Al as long as there is some fine stuff mixed in. It also works if the foamy solid is crushed up before ignition. The delayed boiling can be postponed indefinitely while mixing the stuff by squishing it around in an H2SO4-safe plastic bag with the air squeezed out. Then just scrape it out and after a while in the air it will boil away (this can be sped up by heating with a flame). If there was too much H2SO4 used, it won't ignite.

[Edited on 1-3-2005 by Quince]

Mr. Wizard - 1-3-2005 at 15:26

Did you get it to ignite without putting any sparkler material in it? Are we seeing the reaction between the acid, Al, and sparklers or just Al and acid?

Quince - 1-3-2005 at 17:02

This will not ignite from a propane torch with the fineness of powder I had. I had to use a little piece of sparkler. Once the sparkler bit ignites, it takes about a second for the "thing" to ignite. The sparkler used for starting was about 1/10 by volume, and I just placed the broken bit on top of the "thing". Most of the combusted material from the sparkler piece itself is usually ends up in its own little melted bit, so I can't say that a significant fraction of the sparkler reacted with the "thing". Note that there is no liquid left in the "thing" when I ignite it, so you can't really call it a reaction between acid and Al (and air, as I said reaction 2 in my previous post), as that already has happened. If you look at the pictures, the one before the burning one, you can see a couple of fragments of sparkler on top of the "thing".

I'm sure someone here has Al powder and H2SO4, so please do this experiment. Here's a procedure for those that don't want to read my first post in detail:

- Take a plastic bag that doesn't react (at least for a while) with concentrated H2SO4.

- Put in some Al powder; the particle size does not need to be uniform, but at least some of it should be quite fine, or else it may not ignite.

- Drip in enough concentrated H2SO4 to just wet the powder into a course, crumbly mixture, and squeeze out all air from the bag while squishing the mix around; as H gas is generated and some of the Al dissolves, the mixture will not be now as crumbly and will stick together.

- As long as it's not exposed to the air, it seems to be stable after the reaction finishes; instead, pour it onto your intended crucible, or a brick or something; within a few minutes it will heat up and boil off all or almost all of the liquid (watch out for acid fumes); if that doesn't happen after a while, put a flame to it and it should do it; the result is a foamy dirty-off-grey solid.

- This can be lit either crumbled, or as it is; I find it easiest to light if you flip it over, revealing the bubble holes, and drop in a fragment from a sparkler into one of the bubble holes, so more area is exposed to the sparkler; crumbling the thing tends to get the sparkler to blow the powder all over; now light the sparkler and it should ignite the thing, and you can smell the SO2.

- If it doesn't work, maybe you need more heat from a larger sparkler piece, but usually it means you used too much H2SO4 (or less likely, too little); if you look at the result of the burn, you should see globs of Al, and probably a bit of yellowish tint here and there, which, if you wet, releases some of the rotten egg smell of H2S, so I'm guessing it may be Al2S3 from some side reaction.

[Edited on 2-3-2005 by Quince]

chemoleo - 1-3-2005 at 18:47

First off, could you please reduce this picture with the burning bits (no 6 from the top)? As nice as it is, I have to scroll to the left and right to view all of the text that is below and above, which is annoying. Besides, the 100 kb guard script is likely to kick it out (See Polv's post).

Although nice, I wonder if you expected some miraculous new reaction in a simple system?
You say yourself, you need to light it --> hence the H2 (hydrogen gas liberated by the reaction) ignites, and contributes to the melting of more Al, and a faster reaction between the H2SO4 and Al. So the reaction speeds up until you got a nice self-sustaining reaction, producing nicely burning Al. Al burns just like Mg under the right conditions, it is very much like it. Mg also needs (like Al) a fair amount of heat to ignite in air. H2 will just help the process.
Even if it ignites without an open flame (unlikely), what do you expect other than aluminium sulphates and dehydrated products thereof? Ok, it may be possibled that active Al (freed from the oxide) reacts with the sulphate, forming the sulphide. Dissolve the residue in H2O, if it produces liters of H2S (Al2S3 decomposes in it) then yes you have a secondary reaction between the Al2(SO4)3 and the Al. Only under conditions where this is favourable though, i.e. at high heat. Under those conditions the Al sulphate oxidises free Al presumably. I guess that is interesting and somewhat surprising....

To follow this up, I wonder if anhydrous Al-sulphate reacts with Al powder... according to this it should.


[Edited on 2-3-2005 by chemoleo]

Quince - 1-3-2005 at 20:52

The picture is on my friend's server and I'll reupload a smaller version when he gets home. I didn't realize this was a problem, as I'm used to having my monitor at a high resolution.

I'm not expecting anything miraculous, just an elucidation of the details.

Where is H2 liberated by the reaction? There is H2 produced when the acid is initially added, and it escapes. I don't light it then. A delayed exothermic reaction boils off remaining liquid. That happens when exposed to the air, not when I light it. Only after these two reactions do I light it. This, as well as the Al2S3 and H2S (there is some) etc., I already commented about several times. Please read my posts instead of scanning over the pictures; I know you are probably too busy to spend time on a newbie's posts, but... Anyhow, burning aluminum makes an oxide, whereas I get lots of metallic aluminum (it looks like metal, feels like metal, and conducts like metal, and there are no other metals here, unless it's something melting off the clay pot, which I doubt).

I'll buy some more Al powder and take a closer look at the intermediate stages, but as I'm not a chemist, I was hoping someone knowledgeable here would do the experiment (as clearly my detailed description is not making sense to you guys).

[Edited on 2-3-2005 by Quince]

chemoleo - 2-3-2005 at 04:20

Ok, let me try again:

You have certain uncertainties in your reaction: 1) amount of H2SO4 (%, too), 2) amount of Al 3) Level of oxidation of Al
... and unknowns such as
1) level of impurities in the H2SO4 2) solubility of Al2(SO4)3 in H2SO4

First off, if you really want to characterise the reaction, then weigh out the amounts you use, do it once in stoichiometric proportions (to get solely Al2(SO4)3), and once with the proportions you need to get it to work. Left-over Al suggests that it is sub-stoichiometric, so the reaction of the Al with Al2(SO4)3 is possible in principle (yielding possibly Al2S3 as mentioned above).
Furthermore, how do you know all of the free H2SO4 is converted to the sulphate, before your sparkler ignition? The Al sulphate is likely to be be insoluble in the H2SO4 that it is dispersed in, and thus it is also likely that the Al underneath still hasn't reacted (--> protective coating). Heating, however, may then restart the reaction, in that the protective layer of sulphate is broken up, where the exposed Al can now react - energetically this time, because of the high temps involved (and possible secondary reactions to the sulphide and oxide)

To test this maybe you want to try the following: Take your H2SO4/Al paste, and heat it (small amount) in a test tube, and collect the gasses coming off. Most likely you will get more H2 gas, and possibly decomposition products, such as SO2 and SO3 and H2O.

At last, the fact it's still a crumbly substance before ignition really does suggest not all the H2SO4 has reacted (Theoreticallly it should be a powder). So your 'flash' is probably a result of a number of reactions happening.

[Edited on 2-3-2005 by chemoleo]

Axt - 2-3-2005 at 05:21

Quote:
Originally posted by Bert
Sulfates can be used as oxidizers with Al or Mg to make "flash powder", albeit one with a very high ignition temperature.


I dont know If your still debating the correctness of that, but Bert is entirely correct. Al sulphate with Al more then likely will flare up on ignition, as will it with most other metal sulphates. I know for fact that Mg sulphate with Mg flashes quite quickly.

Quince - 2-3-2005 at 06:11

I know there is no liquid H2SO4 left because the first time I made this I was heating it with a propane torch trying to ignite it -- it wouldn't. However, when I then used the sparkler on the same sample it did ignite. I don't know how hot a propane torch flame is, but I'm pretty sure it's way above the boiling point of H2SO4. I'm interested in the sulfate reacting with the metal, but then why do I have left mostly metallic Al? Unless, as I suggested before, the Al2S3 breaks down due to the high heat, with the S burning off with oxygen from the air (that would explain the SO2 smell); also after burning there is only a bit of Al2S3 (not that much H2S smell when wetted). If so, the only thing left that no one is willing to give me a suggestion for is reaction 2, which occurs ONLY when the wet mixture is left exposed to the air. My only guess is that maybe as the oxide coating is dissolved by the acid and the wet mix gets oxygenated from the air, some of the made-more-reactive-due-to-removal-of-oxide-coating-Al starts oxidizing and that causes the heating and boiling (which, as I said, is sped up by applying heat).

Question: can there be a (possibly multistep) reaction between Al oxide and Al sulfate?

I guess I could try separately making the sulfate, drying it, and then mixing with Al powder and try to ignite. The only problem is that the chemicals won't be in as intimate contact as in the original procedure.

[Edited on 2-3-2005 by Quince]

guaguanco - 3-3-2005 at 09:57

Quote:
Originally posted by Quince
Question: can there be a (possibly multistep) reaction between Al oxide and Al sulfate?

I'd be really surprised. Al2O3 is a very stable chemical.

Quince - 3-3-2005 at 17:05

I haven't been able to do any of the test as I have no Al powder left, and keep getting up too late to get to the pottery store before closing...

Microtek - 4-3-2005 at 01:02

Calcium sulfate hydrates form thermite equivalent mixtures with Al that are actually quite easy to ignite, relative to Fe2O3/Al thermites.
I don't see any reason that Al-sulfate should be less effective. So, my guess would be that the addition of sulfuric acid produces a kind of thermite mixture which you can then ignite ( but as usual, something a little more vigorous than a propane torch is required ). Likely, some of the Al is oxidized by the air instead of the sulfate depending on the coherency of the foam. Oxidizer/fuel imbalances is then responsible for the left over Al.

"Only a fraction of the Al could have oxidized given the amount of Al left, and I don't see how a bit of the Al burning could have made this much heat. "

Well, then I guess you need to re-evaluate your opinion of Al as a fuel; it really is very energetic.

Quince - 4-3-2005 at 01:21

I can make the CaSO4.?H2O from sulfuric acid and lime, but as the ? can be from 0.5 to 2, I'm not sure what ratio of this to Al to use. I'll also try premade Al2(SO4)3. That's assuming when I go to the Mad Potter tomorrow they have usable Al powder.

darkflame89 - 10-3-2005 at 03:09

Regarding this subject, I happen to come across this on the internet: http://yarchive.net/explosives/sulfate.html
Quote:

>>>while experimenting one weekend, I discovered the
>>>consentrated sulphuric acid and Al powder will react exothermically to
>>>form a fluffy grey foam/powder. This powder will then ingite, with no
>>>added oxidized, like flask powder. Raw Al powder will not do this. What
>>>is it. Perhaps some form of ULTRA high surface area Al.
>>
>>Sulfuric acid reacts with aluminum to form hydrogen gas, which is
>>flammable and explosive. The mixture is very hazardous because it
>>can spatter the acid into eyes or on skin. It would be safer and
>>just about as exciting to sit around and watch old tin cans rust.
>>
>>Jerry (Ico)
>
>The production of H2 was why we tried the mixture, but I do not think
>that H2 is the main product. As I said, the reaction yealds a fluffy
>grey powder. Typically we would put a table spoon or two of Al powder in
>a zip lock sandwich bag, then squirt in enough H2SO4 to just wet all the
>Al. We then squeese out all the air in the baggy and squish the mixture
>around untill it is well mixed into an Al-acid mud. If the air is
>squezed out of the bag, the mixture will keep unreacted for many hours,
>but it is NEVER stored mixed. This precaution is just for handling safty
>for the 15 os so seconds it takes to mix and handle. After it is mixed
>the baggy is riped open an thrown into a clay flower pot. The mixture
>will sit for five of ten minutes. It seems that air is needed to
>initiate the reaction. After five to ten minutes the reaction will take
>place. The reaction is very quick and exothermic. Taking about one to
>three seconds, the Al acid mud will expand to about ten times it's
>origional volume, and boil off the remaining sulphuric acid. There may
>well be H2 produced, but it typically will not detonate. You get a cloud
>of sulphuric vapor, but it quickly dissapated (ALWAYS DONE OUTSIDE). As
>I have said you are left with this grey foam. It crumbles when touched,
>but you can still pick up chunks several inches across. This stuff seems
>to be stable. It will not detonate with precussion, and does not seem to
>degrade with time. When ignited the, material in foam form, or
>cruched/crumbled back into a powder form acts like flash powder. A
>brilliand white flash, maby several milliseconds long, with practically
>no residue. Neat stuff, but I never did figure out what it was.
>
>Any ideas?
>
>Wayne S


Sounds like you ended up with a mixture of aluminum sulfate and
aluminum. Aluminum will reduce sulfate quite energetically, and
there are some pyrotechnic and explosive formulas based on the
combination.
The delay in the initial reaction might be due to absorbing moisture
from the air, which accelerates the reaction between concentrated
H2SO4 and Al. Depends on how concentrated your acid was to begin with.

RN

Bert - 10-3-2005 at 07:12

No need to make your own sulfates from scratch... Several of the common ones are cheap OTC. Go down to your local pharmacy. Buy a small container of "Epsom salts" (It's hydrated Magnesium sulfate). Bake the Epsom salts in your oven at 250 F for an hour or so to remove the water. Grind it fine and mix with Al. Ignite. If you really must have Calcium sulfate, go buy some old fashioned plaster of Paris. There is an extensive thread on incendiaries made from plaster and Al on E&W, it does work better if you don't hydrate it first.

Can't get up early enough to make it to a chemical supplier??? Kids these days....Sheesh.

Quince - 11-3-2005 at 20:01

Is this the reaction with the dehydrated Epsom salt?

2 Al + 3 MgSO4 --> Al2O3 + 3 MgO + 3 SO2

Quince - 13-3-2005 at 20:40

I was unable to ignite dried MgSO4 and Al, using a mixture with amounts according to the above equation, even with a sparkler or the Al+H2SO4 thing. The other formula I could think of is 8 Al + 3 MgSO4 --> 4 Al2O3 + 3 MgS, and I may try those ratios.

I'll try KSO4 (fertilizer), as I have lots, but I think the H2SO4 works paritally because of the much more intimate mix that forms. Also, the stuff seems like a simple incendiary to start thermite. Again, I'm not sure what the reaction is.

Edit: adding together stoichiometric amounts of Al powder and H2SO4 to make Al2(SO4)3 does nothing, until heated, and then there's a sudden reaction producing a good deal of smoke. The result is a dark grey powder, rather than the white powder that aluminum sulfate is described as when I searched the Internet. What's going on?

If I use this powder as if it was Al2(SO4)3, then with stoichiometric mixtures testing these two reactions:
2 Al + Al2(SO4)3 --> 2 Al2O3 + 3 SO2
8 Al + Al2(SO4)3 --> 4 Al2O3 + Al2S3
The first doesn't ignite, whereas the second one does, with nothing left in place, but off-white solid residue forming on surrounding surfaces (alumina?); some of the Al2S3 seems to oxidize in air as I can smell SO2 more than H2S. Again, that's assuming that with the sulfuric acid I actually made the sulfate, but then why is it dark grey and give a strange reaction? I read that aluminum sulfate is prepared by reacting alumina, not metallic aluminum, with the acid. Perhaps the reaction only happens as the acid reacts with the oxide coating of the particles, and exposed aluminum gets oxidized by air, and then the acid reacts with that, etc.

With other sulfates, the following seem to work:
8 Al + 3 MgSO4 --> 4 Al2O3 + 3 MgS
8 Al + 3 K2SO4 --> 4 Al2O3 + 3 K2S
Didn't bother trying calcium. So I guess that's that.

[Edited on 14-3-2005 by Quince]

Quince - 13-5-2006 at 03:19

Confirmed with dehydrated copper sulfate. Easier to ignite than CuO-based thermite anyway.

panoptic - 14-5-2006 at 04:18

Stumbled upon this paper the other week, of possible interest.

Attachment: Combustion of mixtures of metal sulfates and magnesium or aluminum.pdf (218kB)
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