Sciencemadness Discussion Board - Aluminium foil to Aluminium Oxide?

Aluminium foil to Aluminium Oxide?

sinai - 17-8-2011 at 16:29

Firstly, I want to know if the product from this is Al2O3 (Aluminium Oxide):

I took a piece of Aluminium foil and lightly crumbled it into a ball. I weighed the Al foil ball, it was 0.334g. I then torched the Al ball with propane until near destruction, the aluminium foil shrunk and changed colour to a more dull grey metallic surface. Once cooled, I weighed the now smaller and harder Al ball and the weight was 0.38g, an increase of 0.044g which I beleive is due to oxidation. Is the product aluminium oxide? Or a partial oxidation?

Secondly, what is a good 1step method to create aluminium oxide from aluminium foil with a high grade yield? I already know of the aluminium and NaOH to create aluminium hydroxide is then heat until it decomposes to aluminium oxide; < which I think is one step too many.

I appreciate any input. And thanks for reading.

not_important - 17-8-2011 at 19:50

Al2O4 is the only stable oxide around STP, so "partial oxidation" would mean having unreacted Al metal. Calculated what weight of Al2O3 you should get with 100% reaction and compare.

And that's the problem - the oxide layer does a good job at protecting the underlying metal that it's difficult to get full conversion. You'd do better finding some Al2(SO4)3 in a garden supply store and precipitating the hydroxide using NaOH, Na2CO2, NaHCO3, or even aqueous ammonia. The product will be purer as well, most metallic Al is some alloy or other.

[Edited on 18-8-2011 by not_important]

condennnsa - 17-8-2011 at 23:10

I've read lots of times on the internet about people making 'aluminum oxide' by heating foil or al cans with a blowtorch, but it just does not work, like not_important said the oxide layer passivates the aluminum very well even at temperatures well above the melting point of Al.
What I think happens is that the metal as it becomes molten flows down from between the two aluminum oxide "shells", this is why after heating you are left with a crumbly thing of what seems to be Al2O3.

Since you mentioned the sodium hydroxide method, I am really fascinated by that reaction, because I once tried it with a very concentrated solution of NaOH and what happened is that no matter how much Aluminum I would add to it, it would keep on reacting , and the liquid became thicker and thicker because of the Al(OH)3 ? slurry. I must have reacted about 400-500g of aluminum in about 50-60g of NaOH.
On the other hand when I used a very dilute solution of NaOH, the reaction would pretty soon stop, and no Al(OH)3 precipitate being formed.

So I think that above a certain pH, the sodium aluminate decomposes to NaOH and Al(OH)3 ??
Is this correct? has any of the more experienced members here done this experiment?

Endimion17 - 18-8-2011 at 00:19

condennnsa is right, any aluminum left is sandwiched between two layers of very thin oxide, and it can not be oxidized further by that method. Aluminum foil burning was always a lousy way of getting the oxide.

There's one interesting thing that shows up when the foil is being heated in different flame zones. Colors appear, similar to ones anodizing produces, and showing that indeed there are oxide layers of different thicknes being produced, although all of them are ridiculously thin compared to the thickness of the foil.

sinai - 18-8-2011 at 01:51

Yeah, I guess the oxide layers just increased, but still a sandwicch.

@not important : yeah I want to buy some, but I've seen brown, white, see through and metallic powders all being called aluminum oxide. Guess some are not as pure as the others. But if there is an easy quick method of making it from Al foil, the price difference would justify it, in the uk anyway.

@condensa : it depends how much al you have, the best way is to use as little water as possible and add the NaOH to the al slowly with water present until a reaction proceeds. And yes PH has something to do with the initiation, but I would say with the dilute NaOH, it was happening very slowly, and perhaps you didn't notice the reaction.

@endimion : yeah.. funny enough, iron oxide and aluminium powder form termite, even though the al powder still has an oxide layer.

Retard-3000 - 18-8-2011 at 02:19

4Al + 3O2 > 2Al2O3

0.33g of Aluminium should oxidize completely to 0.62g of Aluminium Oxide, However, you only have 0.38g of Al/Al2O3 showing that it hasn't been fully oxidized.

blogfast25 - 18-8-2011 at 06:44

Dissolve the foil in 50 % NaOH or KOH (allow a small stoichiometric excess), to form aluminate (NaAl(OH)4), filter solution, if deemed necessary. Simmer to increase reaction speed, if needed.

After cooling, carefully neutralise with about 20 % H2SO4, till close to pH 7. Good pH control is needed to prevent the oxide from redissolving. Gelatinous Al(OH)3.nH2O drops out. Filter and wash filter cake carefully with several small aliquots of clean water. Filtrate contains valuable Na2SO4 (or K2SO4) and can be recovered. Dry the hydrated alumina till constant weight at about 150 - 200 C. Technical Al2O3 is now yours. For certain uses, calcine properly.

[Edited on 18-8-2011 by blogfast25]

sinai - 18-8-2011 at 07:27

@Retard-3000 - yeah, your calculations are correct, I kind of guessed my experiments weight increase was minuscule. But I wanted to double check if it was al oxide, you never know, maybe unburnt propane itself fused with the al :-P Thanks for the calculation.

@blogfast25 thanks, going totry it out. By the way, do you have any idea on how to chemically form al powder?

blogfast25 - 18-8-2011 at 08:35

Quote: Originally posted by sinai

@blogfast25 thanks, going totry it out. By the way, do you have any idea on how to chemically form al powder?

The most technologically advanced method involves turning liquid aluminium into an aerosol (but in argon gas) by pumping it through a suitable nozzle. Good luck with that!

Ball milling appears difficult but possible.

The most affordable method for the amateur is grinding a block of aluminium under water on a turning grinding stone or drum covered with sanding paper. There are designs for such machines published by hobbyists on the Net. Google and yee shall find!

There are no chemical means of doing this, AFAIK.

[Edited on 18-8-2011 by blogfast25]

AJKOER - 18-8-2011 at 09:11

A home chemist preparation based on one of my threads is:

1. Pre-soak Al foil with vinegar for a few hours. This makes subsequent chemical attack on the foil more efficient. I got this from a website where an Aluminum coating manufacturer that lamented over the ability of vinegar to bypass their resistant/hardening efforts on Al foil. Note, if acidic food can easily react with Al, there would be a health issue with Aluminum (assuming there isn't one already).

2. Let treated Al foil slowly dissolve overnight in household ammonia. Depending on your foil, some Silicon (a black solid) will remain.

Actually, the chemistry is that the NH3 dissolves the Al2O3 and then water reacts with the pure exposed Al liberating H2 and forming Al(OH)3, which is gelatinous.

------------------------------
On forming highly reactive Al from solution, I would think one might make a water free Aluminum salt and dissolve in an appropriate organic solvent. Then, react with another metal via a displacement reaction to liberate the Aluminum, may be a path.

[Edited on 18-8-2011 by AJKOER]

[Edited on 18-8-2011 by AJKOER]

Bezaleel - 18-8-2011 at 09:35

Quote: Originally posted by AJKOER

2. Let treated Al foil slowly dissolve overnight in household ammonia. Depending on your foil, some Silicon (a black solid) will remain.

Actually, the chemistry is that the NH3 dissolves the Al2O3 and then water reacts with the pure exposed Al liberating H2 and forming Al(OH)3, which is gelatinous.

If that is really so, then what are the reaction products? I'm amazed to hear aluminium would dissolve in ammonia, as eating through the oxide layer in strong NaOH already takes a few minutes.

condennnsa - 18-8-2011 at 09:51

Quote: Originally posted by Bezaleel

If that is really so, then what are the reaction products? I'm amazed to hear aluminium would dissolve in ammonia, as eating through the oxide layer in strong NaOH already takes a few minutes.

Strong NaOH takes seconds to do that

Another method to make alumina from Al would be by electrolysis, an aluminum anode in NaCl solution will quickly disappear to Al(OH)3 ...

[Edited on 18-8-2011 by condennnsa]

AJKOER - 18-8-2011 at 11:09

Bezaleel:

As was correctly noted by Blogfast25, Aluminum quickly dissolves in NaOH to form Sodium aluminate (formula: NaAl(OH)4 ) and not Al(OH)3. This potentially violent reaction (not referring to Blogfast25's recipe) also produces hydrogen gas. One then must also carefully neutalize (also dangerous) with H2SO4 to form the Al(OH)3 and avoid re-dissolving. These are caustic chemicals that you may or may not have on hand.

On the other hand, with my procedure one can avoid any neutralizing as the NH3 just evaporates. You can also avoid trying to filter the jelly by hand pouring out most of the solution, leaving the gelatinous Aluminum hydroxide and add distilled H2O to wash. When done washing, just leave the Al(OH)3 in an open dish to evaporate.

The hydrolysis of Al in H2O (after the removal/penetration of the Al2O3 with ammonia) proceeds as follows:

2 Al + 6 H2O --> 2 Al(OH)3 + 3 H2

Now as to the dissolving of Al2O3 in NH4OH, there is some controversy as to its description as either Ammonium aluminate (only limited indirect support for its existence) or, more likely, an ammonia peptized Al(OH)3 gel.

I personally save my NaOH (and not to mention H2SO4) for something more interesting like making ferrates.

[Edited on 18-8-2011 by AJKOER]

[Edited on 18-8-2011 by AJKOER]

[Edited on 19-8-2011 by AJKOER]

blogfast25 - 18-8-2011 at 12:02

AJKOER:

STOP PEDDLING RIDICULOUS NONSENSE ON THIS FORUM.

The method I use isn’t only not dangerous (with you everything is dangerous!!!), it works very well. Even with 50 % KOH aluminium foil or cuttings DON’T react ‘violently’, in fact it takes time and some gentle heat. And if you can’t handle caustic soda, please take up PHILATELY, or KNITTING…

Ammonia solution DOES NOT FORM aluminates in significant quantities.

STOP PEDDLING DAMN LIES. :mad:

White Yeti - 18-8-2011 at 12:05

The best way I know to make alumina at a reasonlably large scale is... thermite. Try getting your hands on some thermite and put it into a refractory crucible. If you can, drill some holes into the bottom of the crucible so that some of the molten iron can flow out. Obviously, it's a good idea to put a container underneath to catch the molten metal. Once the reaction mix is cool, grind it into powder by any means necessary. Alumina is a ceramic that will crumble into a powder while iron is a ductile metal that won't crumble, so you can separate iron bits and pieces with a magnet at this point. For extra purity, you can dump the powder into some HCl, this will get rid of both Al and Fe, making them soluble. Once the powder stops bubbling, filter the mixture and dry. You are now the proud owner of pure aluminum oxide :cool:

blogfast25 - 18-8-2011 at 12:18

Quote: Originally posted by White Yeti

The best way I know to make alumina at a reasonlably large scale is... thermite. Try getting your hands on some thermite and put it into a refractory crucible. If you can, drill some holes into the bottom of the crucible so that some of the molten iron can flow out. Obviously, it's a good idea to put a container underneath to catch the molten metal. Once the reaction mix is cool, grind it into powder by any means necessary. Alumina is a ceramic that will crumble into a powder while iron is a ductile metal that won't crumble, so you can separate iron bits and pieces with a magnet at this point. For extra purity, you can dump the powder into some HCl, this will get rid of both Al and Fe, making them soluble. Once the powder stops bubbling, filter the mixture and dry. You are now the proud owner of pure aluminum oxide :cool:

If you're going the thermite route, simple do KClO3 + 2 Al --- KCl + Al2O3. I use this mixture as 'ignition mix' for most pyrometallurgical experiments because it ignites just about any thermite or analog.

That mixture burns so hot that the KCl completely evaporates.

Trouble is that what you obtain is annealed alumina and for most home chemists that's useless: trying to dissolve that is like trying to dissolve corundum: extremely difficult. It's also very hard mechanically speaking: almost impossible to size reduce by hobbyists.

I have a collection of thermite 'left overs' (shells of annealed alumina) and even bashing them with a hammer has almost no effect.

To purify alumina from real thermites is in my experience basically impossible unless you have a way to micronise this extremely hard material...

[Edited on 18-8-2011 by blogfast25]

White Yeti - 18-8-2011 at 12:31

Quote: Originally posted by blogfast25

I didn't think of that, but you're absolutely right, that method would give you a tough mass of annealed aluminium oxide. You could probably use it to make tank armour

I thought the OP wanted to have some alumina powder, not an indestructable lump. Although I'm not sure if thermite would also yield an indestructable annealed lump as well. It's been a while since I set off my last thermite, so I don't remember.

blogfast25 - 18-8-2011 at 13:03

Quote: Originally posted by White Yeti

I didn't think of that, but you're absolutely right, that method would give you a tough mass of annealed aluminium oxide. You could probably use it to make tank armour

In my book (I've 'thermited' oxides of Fe, Si, Ti, Cr, Mn, V, Nb, Cu and some alloys too) a well designed thermite formulation yields good quality metal, well separated from the nuissance by-product, alumina. To achieve this, the mixture has to be designed to reach well above the MP of alumina, otherwise no separation between metal and slag (alumina) can occur. So the alumina, all being well, forms in the molten state (as does the metal) and gravity and immiscibility then cause the metal to separate out and sink to the bottom of the slag puddle. Both metal and slag then solidify, the slag forming annealed alumina. Hard as corundum, trust me. I might post a few photos of these alumina shells here later on.

AJKOER - 18-8-2011 at 13:57

Blogfast25:

As usual you fail to qualify your reputed knowledge.

I have performed this very reaction many years ago using NaOH flakes and Al pieces. The ingedients were directly from a Drano can intended to clear drains. If it was not a voilent reaction in concentrated form, generating both heat and gas as I observed, Drano would probably not work! You should have read the label warning.

Now, I do not necessarily doubt that your NaOH may act differently as we are discussing different concentrations and possibly also quality issues. However, unlike you, I do not assume everyone understands the dangers of working with strong bases and acids, and to critize someone who values safety on a public forum based on experience is inappropriate to say the least.

On the Ammonia aluminate, you have the literature as I gave it to you, and my comment was completely in keeping with it, to quote "only limited indirect support for its existence) or, more likely, an ammonia peptized Al(OH)3 gel"

.

sinai - 18-8-2011 at 15:12

Thanks guys, and hope the heat is gone.

@condennsa, the electrolysis method sounds ideal.

@ajkoer, about some manufacturers being vary about acidic food is right,most foils are mixed with small amounts of silica and other things to minimize the erosion/holes but would entirely depend on what brand/foil type.

@blogfast25 thermite/metals winning over other metals sounds great, I need to get myself some graphite and carbon based casts and experiment, but as White yet said, most products are desired in powdered/high surface area form to be useful in lab scale chemistry. By the way, I was researching the formation of nano sized particles of various metals incl. Al induced electrochemically. Thought I mention it in case it might interest your intellect.

Its funny though, how nature itself produced aluminium oxide in bauxite. The theory is there, but amazing.

And guys, stop arguing, and about safety, anyone who crosses the road without looking both ways first is to blame, not the guy who gave him the directions.

AJKOER - 18-8-2011 at 17:03

Blogfast25:

I have edited my comments on your recipe as it was not my intention to imply that it is per se, a dangerous process.

I was, in fact, thinking of my personal experience with Drano.

Sorry.

Neil - 19-8-2011 at 05:54

Quote: Originally posted by AJKOER

Blogfast25:

As usual you fail to qualify your reputed knowledge.

I have performed this very reaction many years ago using NaOH flakes and Al pieces. The ingedients were directly from a Drano can intended to clear drains. If it was not a voilent reaction in concentrated form, generating both heat and gas as I observed, Drano would probably not work! You should have read the label warning.

Now, I do not necessarily doubt that your NaOH may act differently as we are discussing different concentrations and possibly also quality issues. However, unlike you, I do not assume everyone understands the dangers of working with strong bases and acids, and to critize someone who values safety on a public forum based on experience is inappropriate to say the least.

On the Ammonia aluminate, you have the literature as I gave it to you, and my comment was completely in keeping with it, to quote "only limited indirect support for its existence) or, more likely, an ammonia peptized Al(OH)3 gel"

.

Drano is a mixture of Aluminium, Sodium Nitrate and Sodium Hydroxide. The idea is the aluminium and hydroxide react producing heat and hydrogen which stirs and agitates. The nitrate combines with the hydrogen and converts it to ammonia which reduces the chance of explosion.

I've brought thermite alumina down to ~400mesh powders in an iron mortar and pestle. I've leached it and heated it and swore at it. Never did purify it though.

Al2O3 IS used in tank armour

I've tried burning Al down with fluxes, poor yield.

I've tried going the NaOH route but the effort to product and cost was to high for anything being reagent amounts.

Try looking for local pottery stores, or order it online for pottery.

I've used a electrolysis cell to break it down, it was the fastest and cheapest but venting the hydrogen was a concern. it produces a lot of heat, I fed in scrap Al during the winter and used the heat to warm my room.

I'd also stay away from carbon crucibles if I were you, others may correct me but in my experience the molten Alumina is very effective at producing alumina oxycarbides when given the chance.

blogfast25 - 19-8-2011 at 06:14

This is an example of thermite alumina but not a very good one. This silicon thermite, boosted with chlorate/Al powder, didn’t reach a high enough temperature for the reaction product mix to collect at the bottom of the crucible, so a lot of it froze up on the way down. Despite its porous nature, this little sculpture resists hammering down very well: it's much harder than many a common rock. In it are embedded ‘studs’ of silicon metal (hard to see on the photo), one about a cm wide.

not_important - 19-8-2011 at 06:19

Quote: Originally posted by AJKOER

As was correctly noted by Blogfast25, Aluminum quickly dissolves in NaOH to form Sodium aluminate (formula: NaAl(OH)4 ) and not Al(OH)3. This potentially violent reaction (not referring to Blogfast25's recipe) also produces hydrogen gas. One then must also carefully neutalize (also dangerous) with H2SO4 to form the Al(OH)3 and avoid re-dissolving. These are caustic chemicals that you may or may not have on hand.

On the other hand, with my procedure one can avoid any neutralizing as the NH3 just evaporates. ...

Nope, don't need H2SO4 or HCl, CO2 will do the trick as will NaHCO3. Takes a powerful base to form aluminates, which is why using NH3(aq) doesn't really do the job.

As said above, thermite leftovers are fused alumina - hard, unreactive, and contaminated with whatever metal was being reduced as well as anything else around. Alumina targeted at abrasive use is also the fused variety, it dissolves in HF or fused NaOH. Pottery grade Al2O3 isn't much better, even pottery grade Al(OH)3 can be slow to dissolve in HCl.

Which is why I suggested starting with agricultural/gardening grade Al sulfate and using one of the bases I listed to ppt out nice fresh Al(OH)3.

Bezaleel - 19-8-2011 at 06:20

Quote: Originally posted by AJKOER

On the Ammonia aluminate, you have the literature as I gave it to you, and my comment was completely in keeping with it, to quote "only limited indirect support for its existence) or, more likely, an ammonia peptized Al(OH)3 gel"

Maybe you could post the reference to literature that I may read the article myself? I'm amazed about amonia being capable of reacting with the oxide layer in one way or another.

Neil - 19-8-2011 at 06:21

I used Ti/Mn oxides to produce brittle tiny metal inclusions - it was the easiest to crush slag I managed. Each 5g only took an hour in a very heavy cast iron mortar and pestle.

Other slags were significantly harder.

blogfast25 - 19-8-2011 at 06:22

One route to aluminium salts and alumina that remains worth exploring is Ye Olde Waye of treating clay (kaolin, a mixture of complex silicates/aluminates) with hot conc. sulphuric acid. In ancient times, alum (KAl(SO4)2.12H2O) was mined as a principal source of soluble aluminium (and later name giver to the element). But extraction of aluminium from clays soon took over for the production of ‘synthetic’ alum.

Surely TheWizardisIn could rustle up some medieval recipes for our delectation?

blogfast25 - 19-8-2011 at 07:05

Quote: Originally posted by not_important

Nope, don't need H2SO4 or HCl, CO2 will do the trick as will NaHCO3. Takes a powerful base to form aluminates, which is why using NH3(aq) doesn't really do the job.

Basically a displacement reaction because the carbonate is more stable than the aluminate, right?

2 NaAl(OH)4(aq) + CO2(aq) === > Na2CO3(aq) + Al(OH)3(s) + 4 H2O(l)

Hadn’t thought of that. But it requires bottled CO2 or a CO2 generator…

NaHCO3(s, aq?) + NaAl(OH)4(aq) === > Na2CO3(aq) + Al(OH)3(s) + H2O(l)

But bicar is poorly soluble in water. You just add it as a solid and keep churning?

not_important - 19-8-2011 at 08:17

Some palce it is easy to purchase 'dry ice', which is reasonable handy for immediate use.

Yeah, stir while sprinkling the bicarb in, until fizzing stops. Excess bicarb isn't going to hurt, washing the Al(OH)3 will extract any excess.

Big guys just set the concentration of NaOH right, then heat and add a little fresh Al(OH)3; this gives 9/10 of the Al in solution as Al(OH)3 and restores the NaOH for use in another bauxite extraction.

Clays have much Si and often Fe in them that there is considerable waste of reagents, in these days there's generally easily accessible Al salts to be found = alum itself, the sulfate, chlorohydrate, hydroxide.

AJKOER - 19-8-2011 at 08:45

Some research on Alumina and Ammonia as requested, and also some interesting stuff on AlN which may be present via any thermite reaction given the temperatures involved.

First paper is "The precipitation of aluminium hydrous oxide and its solubility in ammonia" by Prideaux and Henness. The authors noted that the "precipitation by ammonia and its residual solubility should be explicable in terms of the electrochemical properties of the hydroxide and by the theories of the colloidal state, but the position is by no means clear." Also, the authors noted that precipitation from a sulphate solution via alkalis "follows a course which is determined by the amphoteric ionizations of the hydroxide, but is complicated by colloidal phenomena (Britton1)." Further reading introduces even more complicating points as "this is not the isoelectric point of the alumina itself, as the precipitate contains acid radicle." Please note, that the authors use the term "Hydrous Oxides" as defined by H.B. Weiser in his book "Inorganic Colloid Chemistry", Volume II, addressing the properties of Al, Fe and Cr hydroxides that are neither definite hydroxides nor crystal hydrates.

LINK to Full Paper on "The Precipitation of Aluminium Hydrous Oxide and Its Solubility in Ammonia" by Prideaux and Henness

https://docs.google.com/viewer?a=v&q=cache:ydMv90ZG1p8J:...

On the issue of the existence of ammonim aluminate, here is an interesting albeit dated discussion published in "Journal of the American Chemical Society", Volume 38, page 1287 :

"While no such definite evidence of the existence of ammonium aluminate is available, owing to the above mentioned impossibility of securing ammonia solutions of high alkalinity, there seems to be no reason to doubt the analogy of the solutions in ammonia and the fixed alkalies. In this connection, it is interesting to consider the evidence presented by C. Renz (Ber., 36, III, 2751 (1903)). This author dismisses the possibility of the existence of an ammonium aluminate, even though by an indirect method (viz., solution of Al(OH)3 in Ba(OH)2 and subsequent addition of (NH4)2SO4) he was able to obtain a clear solution free from Ba ++ and SO4-, 50 cc. of which contained 0.1 g. Al2O3. The fact, observed by Renz, that freshly precipitated Al(OH)3 is readily soluble in organic amines, far from being an argument against the existence in solution of ammonium aluminate, would appear to indicate that by the solution of aluminium hydroxide in any base, aluminates are formed, the maximum concentration being dependent upon the alkalinity of the resultant solution and its consequent ability to repress the hydrolysis of the aluminate."

http://books.google.com/books?id=FwoSAAAAIAAJ&pg=PA1287&...

Interesting but not necessay definite source as it is a MSDS on Al2O3 (also confirmatory Aluminum MSDS)

Al2O3 is "Slowly soluble in aqueous alkalie solution-forming hydroxides. Very slightly soluble in acid, alkali."
Also, "Very slightly soluble in cold water. Insoluble in hot water."

Assuming complete accuracy of the 1st sentence (we have observed aqueous ammonia slowly dissolve Al2O3/Al and the apparent formation of Al(OH)3 ), my chemical translation of this statement is:

Al2O3 + 3 H2O + 2 OH- --> 2 Al(OH)3 + 2 OH-

Note, I am not stating the formation of a soluble aluminate and further, as written the aqueous alkali (like NH4OH, for example) may act solely as a catalyst. Interestingly, this is precisely the comment on a previously alluded to thread on another forum that was not documented as to source. There is also a parallel to AlN wherein its hydrolysis induction stage is eliminated in the presence of a base with a pH of over 10.

Al2O3 MSDS Source:
https://docs.google.com/viewer?a=v&q=cache:k5ni0DPtE7UJ:...

I also found a similar comment on a chart on page 9 of a report entitled "Aluminum Compounds Review of Toxicological Literature Abridged Final Report", prepared by Integrated Laboratory Systems, namely Al2O3 is "slowly soluble in aqueous alkaline solutions" given possible concerns on the quality of some MSDS statements.

Link:

http://www.scribd.com/doc/2895150/Aluminum

On the burning/thermite reactions with aluminum:

Reference: "Study of aluminum nitride formation by superfine aluminum powder combustion in air" by Alexander Gromov and Vladimir Vereshchagina at Chemical Department, Tomsk Polytechnic University, 30, Lenin Ave., Tomsk, 634050, Russia (available online 18 November 2003).

"ABSTRACT
An experimental study on the combustion of superfine aluminum powders (average particle diameter as0.1 μm) in air is reported. Formation of aluminum nitride during combustion of aluminum in air is focused in this study. Superfine aluminum powders were produced by wire electrical explosion (WEE) method. Such superfine aluminum powder is stable in air but, if ignited, it can burn in self-sustaining way. During the combustion, temperature was measured and actual burning process was recorded by a video camera. SEM, XRD, TG-DTA and chemical analysis were executed on initial powders and final products. It was found that powders, ignited by local heating, burned in two-stage self-propagating regime. The products of the first stage consisted of unreacted aluminum (70 mass%) and amorphous oxides with trace of AlN. After the second stage AlN content exceeded 50 mass% and residual Al content decreased to 10 mass%. A qualitative discussion is given on the probable mechanism of AlN formation in air."

Note, the reaction of any Aluminium Nitride formed in water is reportedly slow with the release of NH3 gas:

AlN + 3H2O --> Al(OH)3 + NH3

Also found a reference on the hydrolysis of AlN which inerestingly parallels that of Aluminum, Title: "The course of the hydrolysis and the reaction kinetics of AlN powder in diluted aqueous suspensions" by Andraž Kocjan, Aleš Dakskoblera, Kristoffer Krnela and Tomaž Kosmača at Engineering Ceramics Department, Jožef Stefan Institute, Jamova 39, SI-1000 Ljubljana, Slovenia
Available online 3 January 2011.

Abstract
"The reactivity of AlN powder in diluted aqueous suspensions in the temperature range 22–90 °C was investigated in order to better understand and control the process of hydrolysis. The hydrolysis exhibits three interdependent stages: during the induction period (first stage) amorphous aluminum hydroxide gel is formed, followed by the crystallization of boehmite (second stage) and bayerite (third stage). The hydrolysis rate significantly increased with higher starting temperatures of the suspension, but was independent of the starting pH value; however, the pH value of 10 caused the disappearance of the induction period. The kinetics was described using un-reacted-core model, and the chemical reaction at the product-layer/un-reacted-core interface was the rate-controlling step for the second stage of the hydrolysis in the temperature range 22–70 °C, for which the calculated activation energy is 101 kJ/mol; whereas at 90 °C, the diffusion through the product layer became the rate-controlling step."

sinai - 19-8-2011 at 09:06

@blogfast25 that piece sure looks nice :-) yeah, al oxide is v. Hard, hence it being used on sand papers.

@Neil yeah, I always redirect heat into the house so to not waste the energy. I also tried the NaOH method, and it isn't too economic, right now i'm trying the electrolysis methods, experimenting with ways to accumalate a powder form. E.g. pulsing the electrolysis.

Right now im waiting on a decision for Kingstn University for Pharmaceutical sc. and funnily the faculty leader is also called Neil. Hope I get in. So guys pray for me

Neil - 19-8-2011 at 11:12

Do you mean forming powdered oxide or metal?

I used a fixed DC current and ended up with slime which would turn into a wet powder as more aluminum was fed in. I set up two ring shaped electrodes which I set aluminum extrusion in so the whole thing was self feeding.

@AJKOER - wtfbbq? You're spamming with irrelevant gibberish...

blogfast25 - 19-8-2011 at 12:00

Quote: Originally posted by Neil

@AJKOER - wtfbbq? You're spamming with irrelevant gibberish...

It's largely a replica of the previous run: reports from 1916, MSDSs etc etc. Some things never change...

White Yeti - 20-8-2011 at 01:08

Reaction of aluminium chloride with sodium hydroxide will yield an aluminium hydroxide precipitate. Filter, dry, and dehydrate to get aluminium oxide.

Guys, this not a tough question STOP ARGUING! Every day I spend on this forum, people are beating each other up, getting angry at one another, arguing etc... Learn to get along, science is not what it used to be. To get something done, you can't work alone, you need the help of others, you need feedback, input, fresh and original ideas and information, cooperation and constructive criticizm. I thought this forum would bring intelligent people together, promote cooperation, but no. All you guys do is argue amongst one another, trying to prove you are smarter and better than the rest, shrinking those who are already small. Nothing gets done when you argue, and time is wasted.

Bezaleel - 20-8-2011 at 04:32

Quote: Originally posted by blogfast25

It's largely a replica of the previous run: reports from 1916, MSDSs etc etc. Some things never change...

blogfast25, it was me who asked for the information, the "previous run". Edit: and I'm happy it was posted.

In addition, in a few respects, reports from the beginning of the 21st century are more valuable to a simple home chemist than that which is published these days. Simply for reasons like a home chemist doesn't usually work with ultra pure chemicals, usually has no access to advanced measurement methods like chromatography, X-ray diffraction, etc. And so, information such as a description of colour, appearance, and ease of dissolution are of much practical use to a simple chemist. Your experiences may be different, but mine are that such information is more frequently found in the older publications.

[Edited on 20-8-2011 by Bezaleel]

blogfast25 - 20-8-2011 at 05:07

Quote: Originally posted by Bezaleel

In addition, in a few respects, reports from the beginning of the 21st century are more valuable to a simple home chemist than that which is published these days. Simply for reasons like a home chemist doesn't usually work with ultra pure chemicals, usually has no access to advanced measurement methods like chromatography, X-ray diffraction, etc. And so, information such as a description of colour, appearance, and ease of dissolution are of much practical use to a simple chemist. Your experiences may be different, but mine are that such information is more frequently found in the older publications.

[Edited on 20-8-2011 by Bezaleel]

I’m not opposed to using older reports and do it all the time. But often the information has been superseded, by use of HiTech or not (it matters not one jot - being able to understand modern texts does not rely on being able to replicate the experimental data contained therein).

AJKOER’s claims of being able to dissolve aluminium with ammonia flies in the face of the simple fact that ammonia is used frequently in analytical chemistry to precipitate hydrated alumina from aluminium salt solutions because the weakly alkaline ammonia solution doesn’t have the power to dissolve ANY alumina. By contrast, using stronger alkalis like alkali metal hydroxides or alkali metal carbonates, risks ‘overshooting’: add a bit too much and alumina starts to redissolve again because it is amphoteric. These observations are perfectly in agreement with theory. Rarely does one see such an open and shut case.

AJKOER’s observations, by contrast, remain largely shrouded in mystery. We’re not explained what exactly he’s done, which reagents and at what strength were used, in what conditions. No calculations or quantitative data were ever presented.

Well, if you’re going to make strong and unusual claims, present strong evidence. If not, others will have the tar and feathers at the ready. T’was thus in olden time, t’is thus still today.

Neil - 20-8-2011 at 05:46

Quote: Originally posted by White Yeti

...Guys, this not a tough question STOP ARGUING! Every day I spend on this forum, people are beating each other up, getting angry at one another, arguing etc... Learn to get along, science is not what it used to be...

No argument there.

AJKOER is posting irrelevant pieces of data out of context. For example he says that the nitride data is in relation to thermites, it has nothing to do with thermite.

Further, by selectively quoting the article he omits the data that the aluminum was not burned in open air but rather in a sealed reactor - the difference is night and day. Aluminum nitride has nothing to do with this thread.

That is not bad science, that is fraud. Up until that the thread was largely in self agreement. Pissing contests are one thing and yes, they bring the quality of the forum down. Posting BS and then re-posting it over and over and over again should be met with rigid objection because this is science, not BS.

AJKOER - 20-8-2011 at 17:59

I apologize for not being even more clear and adding the person's name as to why I was supplying the references.

OK, to get the record straight, I believe the reaction of ammonia and Al2O3 is best described as the ammonia acting solely as a catalyst. The mechanics, I suspect, are that the ammonia raises the pH and reduces the inception period as has been observed in the hydrolysis of AlN. Thus, the product of NH3 and Al2O3 is most likely a NH3 peptised/supersaturated hydrated alumina (with no chemical reaction). Other possibilities are less likely, in my opinion. On a point of full disclosure, my catalyst depiction is not an original opinion, as it was presented on another thread, but left undocumented as to source (in spite of my direct attempts to secure documentation).

However, interesting from a patent law perspective, since one might claim that there is an ostensible reaction that seemingly produces Ammonium aluminate (as I described in my dated reference from Journal of the American Chemical Society, a legally authentic source), one cannot, from a patent law perspective, categorically declare that is does not exist. Hence, interestingly I have seem references to it in Patents, possibly to cover bases, and in my opinion, not due to its likely existence. This legal point is also why I find it important to quote the reference, however foreign to my personal opinion. My source is taken from the US Government Patent Office ( per 2173.05(t) Chemical Formula - 2100 Patentability), "A compound may also be claimed in terms of the process by which it is made without raising an issue of indefiniteness", where indefiniteness is a cause to dismiss a patent claim. Interestingly, Ammonium aluminate has been cited in at least one patent, possibly since one cannot claim its "indefiniteness".

Now, my presentation of the AlN material was mostly due to its interesting hydrolysis properties as described earlier and to confirm the observed difficulties in burning Al to form Al2O3. I also noted that it MAY (or may not especially in completely sealed environments) be of any significance in thermite reactions. However, to be honest, when people are producing some very hard materials at home, given the hardness properties of AlN, I wonder if any AlN could be present in spite of claims of purity from home chemists (which, may indeed, be the case here).

GOOD LINK TO PRIDEAUX PAPER
https://docs.google.com/viewer?a=v&q=cache:ydMv90ZG1p8J:...
[Edited on 21-8-2011 by AJKOER]

[Edited on 21-8-2011 by AJKOER]

AJKOER - 21-8-2011 at 01:43

On references, I should also cite an influential one on my thinking supplied by Neil. To quote Neil's specific reference taken from : "Pitting Corrosion Mechanisms and Characterization of Aluminum in solar Heating Systems" by Liang and Zhang (2006):

"The complex characteristics and mechanisms of aluminum pitting corrosion in a solar heating system were studied by the chemical immersion method and electrochemical techniques as well as fractal theory. The results showed that pitting corrosion of Al occurred in a tap water environment due to the local enrichment of Cl- ions."

My favorite quote from this paper is:

"However, the passivation film over Aluminum and Al alloys is easily destroyed in the solution containing active anions, such as Cl-, which leads to localized corrosion".

I believe this is directly referring to the removal of the protective Al2O3 layer by Chloride anions. My take on this is that there may be many compounds that are possible catalysts to induce the Al hydrolysis reaction by disrupting the pH of the solution for which the authors noted that "Aluminum and Al alloys are passivated in neutral solutions".

LINK:
http://proj3.sinica.edu.tw/~chem/servxx6/files/paper_7425_12...

I would also mention that Neil performed experiments seemingly confirming the action of chloride anions on Al as the paper noted.

[Edited on 21-8-2011 by AJKOER]

AJKOER - 21-8-2011 at 02:23

A point on a common thermite starting mixture:

2 KNO3 + 4 Al + S --> K2S + N2 + 2 Al2O3

So, it is possible to have nitrogen gas present even in a sealed environment, although this may not be the case here.

More of concern to me does any of the very hot Al2O3/Al mix cool in air, and could this produce any AlN? Note the surprising large percent of the arc heated Al converted to AlN in my cited reference when burned in open air.

However, if any of you more experienced thermite guys are convinced that AlN is not a factor, I will accept it.

[Edited on 21-8-2011 by AJKOER]

blogfast25 - 21-8-2011 at 04:46

Quote: Originally posted by AJKOER

A point on a common thermite starting mixture:

2 KNO3 + 4 Al + S --> K2S + N2 + 2 Al2O3

So, it is possible to have nitrogen gas present even in a sealed environment, although this may not be the case here.

More of concern to me does any of the very hot Al2O3/Al mix cool in air, and could this produce any AlN? Note the surprising large percent of the arc heated Al converted to AlN in my cited reference when burned in open air.

However, if any of you more experienced thermite guys are convinced that AlN is not a factor, I will accept it.

[Edited on 21-8-2011 by AJKOER]

Sulphur is sometimes added to these ignition mixes but it's not strictly speaking necessary. The sulphur would combine with the Al though, because Al2S3 has a high value for HoF. That's why the system 2 Al + 3 S === > Al2S3 is used to heat boost silicon thermirtes. Smelly but effective and cheap...

The overwhelming majority of thermites aren't carried out in 'sealed environment', as that would be a general call for explosions to occur.

AlN in thermite conditions? Forget it. Firstly AlN has a much lower HoF than Al2O3 which means that the oxide is much, much more preferentially formed than the nitride.

Also: here the oxidising power is provided by the oxydiser (the oxyde, mixture of oxydes or oxyde + booster oxidiser). Nitrogen simply doesn't get a look in.

[Edited on 21-8-2011 by blogfast25]

Bezaleel - 28-8-2011 at 07:07

8 days ago, I put some strong ammonia (18/20° Baume) in an erlenmeyer flask and put a piece of aluminium foil in it. Teeny tiny amounts of gas evolved, the bubbles being on the verge of visibility for the human eye. This was left to react for 8 days, at a temperature of around 20°C.

Today, I took out what has remained of the foil, and poured off the solution. On the flask, a white substance had formed. It did not dissolve in 10% acetic acid, and neither in 5% hydrochloric acid, not even when heated to boiling.

Below are some pictures of the results. The substance formed was pure white, but in the pictures it is a bit coloured due to the lighting conditions.

blogfast25 - 28-8-2011 at 07:36

Try dissolving it in stronger acid and stronger alkali (NaOH or KOH). If it dissolves in the latter you can be fairly sure it's (hydrated) alumina.

If it really is hydrated alumina, what this would clearly suggest is that there's enough OH- (but not much: even at that strength an ammonia solution is a weak base) to attack Al and produce hydrated alumina but not enough to create aluminate (Al(OH)4-

because:

NH3(aq) + H2O(l) < === > NH4+(aq) + OH-(aq)

... leans strongly to the right.

[Edited on 28-8-2011 by blogfast25]

condennnsa - 28-8-2011 at 07:45

shouldn't aluminum hydroxide dissolve readily in dilute HCl?

blogfast25 - 28-8-2011 at 08:03

Quote: Originally posted by condennnsa

shouldn't aluminum hydroxide dissolve readily in dilute HCl?

Depends a bit on strength and age of the oxide, really... Of course the precipitate could have been something in the ammonia but I can't think what it could be.

[Edited on 28-8-2011 by blogfast25]

Neil - 28-8-2011 at 09:32

Crystallized oxide dropped when the material under it was dissolved?

Possibly protective coatings?

blogfast25 - 28-8-2011 at 12:43

Quote: Originally posted by Neil

Crystallized oxide dropped when the material under it was dissolved?

Possibly protective coatings?

Too much stuff for that, IMHO...

Bezaleel - 28-8-2011 at 16:40

Quote: Originally posted by blogfast25

Try dissolving it in stronger acid and stronger alkali (NaOH or KOH). If it dissolves in the latter you can be fairly sure it's (hydrated) alumina.

If it really is hydrated alumina, what this would clearly suggest is that there's enough OH- (but not much: even at that strength an ammonia solution is a weak base) to attack Al and produce hydrated alumina but not enough to create aluminate (Al(OH)4-

because:

NH3(aq) + H2O(l) < === > NH4+(aq) + OH-(aq)

... leans strongly to the right.

[Edited on 28-8-2011 by blogfast25]

To the left, that is... And indeed, no aluminate is formed.

Since not all of the foil dissolved, the white material is not the oxide skin of the foil.

IIRC, aluminium hydroxide is a voluminous, jelly substance, so I'm inclined to say the white substance be the oxide. In addition, the hydroxide would have easily dissolved in 5% HCl.

I discarded the reaction products, so I can't do any furhter testing. I'm amazed though, that with a weak base like ammonia, the foil has reacted at all.

Neil - 28-8-2011 at 19:46

In the first pictures it looks like there is more white material, a slime or gel?

Is it possible that some of the material was lost when you poured of the original liquid or during subsequent filtration/decantations?

I agree, the amount of material is to great. Ammonia is used in a number of anodizing mixtures for Al, is it possible it catylised the crystallization of the aluminum hydroxide?

The oxide layer on the foil may have been thick enough to provide seed crystals.

Did you happen to handle the foil with your bare hands before it was added to the ammonia?

AJKOER - 28-8-2011 at 20:24

Caution, while the voluminous jelly substance is most definitely Al(OH)3, the white substance you saw could be an impurity from the Aluminum foil as it is commercial grade Al, which is anywhere from 92% to 99% pure, or some alpha Al2O3, formed in a hardening process, which is not soluble in acids or bases.

On impurities, we have mentioned Silicon, some foils have Fe for added strength, at times Mn, but would anyone believe a small amount of Lead (Pb maybe country specific due to manufacturing process) as a possible impurity? See article reference below (I hope I am misreading).

"The effect of lead impurity on the DC-etching behaviour of aluminum foil for electrolytic capacitor usage" by W. Lina, , G. C. Tua, C. F. Linb and Y. M. Pengb
a Institute of Materials Science and Engineering, National Chiao Tung University, Hsinchu, Taiwan R.O.C., at Materials Research Laboratories, Industrial Technology Research Institute, Hsinchu, Taiwan R.O.C.
Received 26 June 1995; revised 17 October 1995. Available online 16 February 1999.
"Abstract
The effects of lead impurity on the etched morphology of high purity aluminum foils for electrolytic capacitor applications were investigated in this work. The lead impurity was either present in as-received aluminum foils or deposited purposely on the foil surface through an immersion-reduction reaction. The amount and distribution of deposited lead varies with the lead content in as-received foil. The as-received foil with higher lead content gave a higher concentration and a more uniform distribution of deposited lead."

Note, Aluminium is reported to be chemically resistant in contact with substances in the pH range 4 to 9.

Neil - 28-8-2011 at 20:53

Quote: Originally posted by AJKOER

Caution, while the voluminous jelly substance is most definitely Al(OH)3, the white substance you saw could be an impurity from the Aluminum foil as it is commercial grade Al, which is anywhere from 92% to 99% pure, or some alpha Al2O3, formed in a hardening process, which is not soluble in acids or bases.

On impurities, we have mentioned Silicon, some foils have Fe for added strength, at times Mn, but would anyone believe a small amount of Lead (Pb maybe country specific due to manufacturing process) as a possible impurity? See article reference below (I hope I am misreading).

"The effect of lead impurity on the DC-etching behaviour of aluminum foil for electrolytic capacitor usage" by W. Lina, , G. C. Tua, C. F. Linb and Y. M. Pengb
a Institute of Materials Science and Engineering, National Chiao Tung University, Hsinchu, Taiwan R.O.C., at Materials Research Laboratories, Industrial Technology Research Institute, Hsinchu, Taiwan R.O.C.
Received 26 June 1995; revised 17 October 1995. Available online 16 February 1999.
"Abstract
The effects of lead impurity on the etched morphology of high purity aluminum foils for electrolytic capacitor applications were investigated in this work. The lead impurity was either present in as-received aluminum foils or deposited purposely on the foil surface through an immersion-reduction reaction. The amount and distribution of deposited lead varies with the lead content in as-received foil. The as-received foil with higher lead content gave a higher concentration and a more uniform distribution of deposited lead."
...

AJKOER - 28-8-2011 at 22:12

Thanks Neil for the highlights.

I probably misread and now would say there maybe special purpose Al foils with added Pb, for example, for specific applications.

IrC - 29-8-2011 at 00:05

http://www.exo.net/~pauld/activities/AlAirBattery/alairbatte...

Using salt water has already been mentioned, but you could also power something while waiting for the reaction.

blogfast25 - 29-8-2011 at 03:59

Quote: Originally posted by AJKOER

Thanks Neil for the highlights.

I probably misread and now would say there maybe special purpose Al foils with added Pb, for example, for specific applications.

Lead? In what is in all likelihood food grade Al foil? I think NOT!

Bezaleel - 29-8-2011 at 16:21

Quote: Originally posted by blogfast25

(...)

Lead? In what is in all likelihood food grade Al foil? I think NOT!

In all certainty food grade Al foil.

Quote: Originally posted by Neil

In the first pictures it looks like there is more white material, a slime or gel?

Is it possible that some of the material was lost when you poured of the original liquid or during subsequent filtration/decantations?

I agree, the amount of material is to great. Ammonia is used in a number of anodizing mixtures for Al, is it possible it catylised the crystallization of the aluminum hydroxide?

The oxide layer on the foil may have been thick enough to provide seed crystals.

Did you happen to handle the foil with your bare hands before it was added to the ammonia?

No material was lost when I poured off the liquid, as it stuck a bit to the walls of the flask. It was a very thin layer, that broke into many pieces, when I shook the flash after I added the HCl solution. In the 3rd pic you see the broken pieces of what grew on the flask walls in the 1st pic.

This is of course food grade Al foil, and I tore off a piece from the role by hand (so there are human lipids on it for sure).

The original question in this thread was a method to produce pure Al2O3. It seems, putting pieces of household Al foil in ammonia is a possible though slow method. An interesting test might me to see whether the reaction speeds up when refluxing the ammonia.

Neil - 29-8-2011 at 18:15

Toss in a chloride to speed it up.

It would be interesting to see if it reacts without any salt/oil contamination from ones hands.

Do you have the capacity to reflux it in ammonia? That would be very interesting...

[Edited on 30-8-2011 by Neil]

blogfast25 - 30-8-2011 at 04:09

Quote: Originally posted by Bezaleel

The original question in this thread was a method to produce pure Al2O3. It seems, putting pieces of household Al foil in ammonia is a possible though slow method. An interesting test might me to see whether the reaction speeds up when refluxing the ammonia.

Considering just how many quick and dirty methods there are to produce alumina this would definitely be a slow boat to China. And you still need to prove the stuff you see is actually alumina.

An interesting but rarely discussed laboratory method for producing calcined alumina is the pyrolysis of pure ammonium alum (NH4Al(SO4)2.12H2O), which on heating first loses its water, then the ammonium sulphate, then the SO3.