Sciencemadness Discussion Board

Reaction between NH4OH and Al2O3

AJKOER - 13-6-2011 at 08:01

I conducted a home experiment of first burning Aluminium foil till red hot (forming brittle Al2O3) and adding the freshly prepared Al2O3 to NH4OH. The reaction produces tiny gas bubbles (H2?) and some evident of gelatinous like Al(OH)3. So what is the precise reaction?

Candidate 1: The NH3 is being decomposed per a reaction between NH3 and Al2O3 as it does with many metal oxides (MxO):

NH3 + MxO --> H2O + N (or on occasion NO) + M (or on occasion MN)

In this case with the Aluminium oxide and NH4OH, albeit not necessarily a rapid reaction:

2 NH4OH + Al2O3 --> 5 H2O + N2 + 2 Al

and immediately:

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

So, the combined reaction is:

2 NH4OH + Al2O3 + H2O --> 2 Al(OH)3 + N2 + 3 H2

and with excess ammonia, the possible creation of the Al(OH)4-, given the amphoteric nature of Aluminium:

3 NH4OH + Al2O3 + H2O --> 2 NH4[Al(OH)4] + N2 + 3 H2

Perhaps unlikely but there in the literature some support with, for example, an article that examines the low temperature oxidation of NH3 with a Ag/Al2O3 catalyst. Reference : "Mechanism of selective catalytic oxidation of ammonia to nitrogen over Ag/Al2O3" by Li Zhang, Hong He, based on their work at the State Key Laboratory of Environmental Chemistry and Ecotoxicology, Research Center for Eco-Environmental Sciences, Chinese Academy of Sciences, Beijing 100085, PR China.

Here is the abstract:

"The mechanism of selective catalytic oxidation (SCO) of NH3 over Ag/Al2O3 was studied by NH3 temperature-programed oxidation, O2-pulse adsorption, and in situ DRIFTS of NH3 adsorption and oxidation. The essence which affects the low temperature activity of Ag/Al2O3 has been elucidated through the mechanism study. Different Ag species on Ag/Al2O3 significantly influence O2 uptake by catalysts; while different oxygen species affect the activity of NH3 oxidation at low temperature. The activated –NH could react with the atomic oxygen (O) at low temperatures (<140 C); however, the –NH could also interact with the O2 at temperatures above 140 C. At low temperatures (<140 C), NH3 oxidation follows the –NH mechanism. However, at temperatures above 140 C, NH3 oxidation follows an in situ selective catalytic reduction (iSCR) mechanism (two-step formation of N2 via the reduction of an in situ-produced NOx species by a NHx species)."

Apparently the reaction is evident at 140 C and my proceed much more slowly at lower temperature.

There is also at least one old chemistry text reports that by shaking fine copper filings (far less reactive then Al) in NH4OH, a white cloud is visible at room temperature, even though Cu (and CuO) catalyst to decompose NH3 now normally operate at much higher temperatures to achieve efficient yields.

This link gives the full free text article per an apparent policy of the People Republic of China to distribute environmental related studies for no charge (thanks).

http://hehong.rcees.ac.cn/bookpic/20101301643540128.pdf

CANDIDATE 2: the creation of Al(OH)4 complex and the presence of some Al in the Al2O3, the freed Al then reacting with water to liberate H2:

2 NH4OH + Al2O3.Al + 3 H2O --> 2 NH4[Al(OH)4] + Al

and immediately:
Al + 3 H2O ---> Al(OH)3 + 3/2 H2

I do not have a source to support the presence of the Al/Al2O3 amalgam. Note, if pure Al2O3 still reacts with NH4OH to produce H2, then this candidate is eliminated.

Other candidates welcome.

[Edited on 13-6-2011 by AJKOER]

AJKOER - 14-6-2011 at 06:34

I believe I found the answer.

The correct answer is, in fact, Candidate 2:

2 NH4OH + Al2O3.Al + 3 H2O --> 2 NH4[Al(OH)4] + Al

and immediately:

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

as it appears that the complete combustion of Aluminium, even as a fine powder, is difficult to accomplish. Per the reference below, the final combustion of Aluminium results in Al2O3, Al (around 10%) and small amounts of AlN (Aluminium Nitride).

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, for those seeking a reference for the reaction of some amphoteric metal oxides (like Al, Cr and Zn) in water see the extract below from http://www.angelfire.com/theforce2/tutorboard/AB12.html

"Amphoteric Metal Oxides

Al2O3(s) + 2OH^-(aq) + 3H2O(l) -> 2[Al(OH)4]^-(aq)

Cr2O3(s) + 2OH^- (aq) + 3H2O(l) -> 2[Cr(OH)4]^-(aq)

ZnO(s) + 2OH^-(aq) + H2O(l) -> [Zn(OH)4]^2-(aq)"

The lesson to be learned here is that the combustion of certain metals in air is a generally impure way to obtained the metal oxides, and in the special case of Al2O3, adding NH4OH is one way to test for the presence of elemental Aluminium via a gaseous reaction mechanism.

bbartlog - 14-6-2011 at 12:03

If your proposal regarding the initial reaction is correct (...'2 NH4OH + Al2O3.Al + 3 H2O --> 2 NH4[Al(OH)4] + Al'...), then adding NH4OH would *not* be a useful test for elemental aluminum in the oxide, as some of the aluminum would be getting regenerated (reduced) even if you started with pure Al2O3. Now, as it happens, I very much doubt that candidate 2 is correct (Al2O3 is very hard to reduce to elemental aluminum and I do not believe that mere NH4OH can do it, even as some sort of intermediate step). So your conclusion regarding NH4OH as a way to test for residual unreacted aluminum is quite possibly correct and sounds interesting, but is at variance with your earlier reasoning.

blogfast25 - 14-6-2011 at 13:19

AJKOER:

I won’t mince my words. Instead of writing pretty, balanced equations you should get a decent but basic textbook on chemistry and study it. You write out reactions which almost without exception are THERMODYNAMICALLY UNFAVOURABLE, hence impossible except for strong forcing conditions. See this set for example, all of which simply CANNOT PROCEED, because the change in Free Gibbs Energy is positive:

2 NH4OH + Al2O3 --> 5 H2O + N2 + 2 Al
2 NH4OH + Al2O3 + H2O --> 2 Al(OH)3 + N2 + 3 H2
3 NH4OH + Al2O3 + H2O --> 2 NH4[Al(OH)4] + N2 + 3 H2
2 NH4OH + Al2O3.Al + 3 H2O --> 2 NH4[Al(OH)4] + Al

Next you quote ill-digested sources to back your wild theories.

Or take this gem:
Quote: Originally posted by AJKOER  
I conducted a home experiment of first burning Aluminium foil till red hot (forming brittle Al2O3) and adding the freshly prepared Al2O3 to NH4OH. The reaction produces tiny gas bubbles (H2?) and some evident of gelatinous like Al(OH)3. So what is the precise reaction?


You provide no evidence WHATSOEVER that the observed ‘reaction’ produces gas, let alone that it’s hydrogen, yet using the ‘observations’you conclude burning metals to oxides is difficult.

What you saw in all likelihood is air escaping from the porous oxide. PERIOD!

LanthanumK - 14-6-2011 at 16:58

Only the Wikipedia science reference desk and similar forums designed to provide reliable help to posters are supposed to harshly reprove promoters of wild chemical theories I was and AJKOER is. Here, it is a discussion forum, and people may propose wrong ideas to get community insight on them. It would be easy to state scientifically why these equations cannot happen under normal conditions without !'s and CAPS.

blogfast25 - 15-6-2011 at 04:19

Quote: Originally posted by LanthanumK  
It would be easy to state scientifically why these equations cannot happen under normal conditions without !'s and CAPS.


… which I have done (learn to read). The majority of these equations represent reactions that cannot proceed because the change in Gibbs Free Energy between left and right is (strongly) positive. AS already STATED. Perhaps caps are necessary in your case too?

An undergrad student of chemistry knows that an ammonia solution is useful for precipitating Al(OH)<sub>3</sub> hydrate from an Al<sup>3+</sup> solution, NOT for dissolving it as aluminate. The pOH of even a concentrated solution of ammonia is too high for the aluminate formation to take place.

Simples.

LanthanumK - 15-6-2011 at 04:36

You did state it scientifically, but with overemphasis.

blogfast25 - 15-6-2011 at 04:51

Sometimes overemphasis is good.

AJKOER - 18-6-2011 at 08:08

My bias is an experimentalist as there are many reactions (especially with regard to catalyst) for which you would read that honest scientist admit that, as of yet, they do not understand the reaction path. Or, in your language, the current reaction is "THERMODYNAMICALLY UNFAVOURABLE", albeit, the reaction does somehow presumably occurs.

Thanks for your opinion in largely discrediting Candidate I, which I, without calculations, was largely suspicious of, although I thought it was very interesting that a Ag/Al2O3 catalyst was able to produce a decomposition reaction on NH3 at 140 C in a gaseous state. As you may be aware of, ferrates can decompose aqueous NH3 at room temperature, but no reason to believe that they were at work here.

However, you did include the Candidate 2, in your no fly list. Do this mean you have a better path than I suggested?

Do you doubt the published research I cited (which came up under a search for Aluminum Nitride)?

If you have questions on the experiment, please repeat it yourself (thoroughly burn pure Al (and/or Aluminum foil) in air (or pure O2), react with pure and/or household ammonia. Does all the compound, with the exception of Si residue in the case of Al foil, eventually dissolve with the formation of a gas?

I performed this reaction many times and also place the reactants in a partially crushed plastic bottle, which under pressure from the gas generated, popped back into shape so be careful. One time the reaction was particularly interesting as the final reactant liquid was clear (excess NH4OH?) and I wanted to react the solution (NH4OH+Al(OH)3=?) with something else (it was MgSO4, very cool, forms a double precipitate Ammonium Aluminum Sulphate and Mg(OH)2 ), so I poured about half of the solution out leaving some still solid Al/Al2O3 in the bottle. The next day, the one time clear solution was cloudy with obvious jelly like Al(OH)3! I only poured some out! What I also found interesting is that some authors mention that excess NH4OH does produce a clear solution with Al(OH)3 and others, including yourself, say it doesn't. To quote: "This precipitate of Al(OH)3, which is amphoteric, dissolves in an excess of hydroxide or in acids". Source: http://www.public.asu.edu/~jpbirk/qual/qualanal/aluminum.htm...

The source could be trustworthy as it is basically pictures of actual reactions. I personally am in the "it depends camp" meaning is there still Al or Al2O3 to be reacted with, pH, etc.

Thanks, however, for being skeptical, for this is how I believe knowledge progresses!

[Edited on 18-6-2011 by AJKOER]

[Edited on 18-6-2011 by AJKOER]

m1tanker78 - 18-6-2011 at 08:23

Ajoker: What method are you using to burn the Al foil? Seems easy enough to try and reproduce...

Tank



sternman318 - 18-6-2011 at 08:56

Well I think you can do a few things to investigate

Try adding the Al2O3 to a base such as NaOH to see if the basic nature of NH4OH causes the reaction to take place

Use a splint test to try and determine if the gas is H2

Have the sample of Al2O3 sit in some water first, then add your ammonia. This will address the possibility of bubbles of air in the oxide sample

try wafting some of the gas to determine it is releasing ammonia ( this is unreliable, as the solution should already smell like ammonia...

Please tell me you are not adding the hot sample straight to the solution?

AJKOER - 18-6-2011 at 09:11

This was actually a home experiment. With caution, if you displace slightly those metal lids on your gas stove, you may notice a strong gas blast in a particular direction (normal methane flame is about 900 C, but with turbulence perhaps as high as 1250 C). I burned cheap Aluminum foil till it glowed red and then, repeated heating the Al2O3 till it glowed again. The Al2O3, however, does not retain heat and you can also immediately handle it. Within 5 minutes of preparation, the Al2O3/AlN/Al was added to NH4OH.

All aqueous reactions proceeded at room temperature and within a few minutes there is obvious evident of a reaction. generating small gas bubbles and Al(OH)3, which continues unabated for hours (so this is not trapped air). Apparently, however, there is still some unreacted Al (and AlN) along with the Al2O3, as one possible explanation.

To answer a question, I do not possess any NaOH that isn't heavy contaminated with other compounds (including NaClO).


[Edited on 18-6-2011 by AJKOER]

[Edited on 18-6-2011 by AJKOER]

LanthanumK - 19-6-2011 at 03:02

NaOH is not all that hard to make.

redox - 19-6-2011 at 05:13

Quote: Originally posted by LanthanumK  
NaOH is not all that hard to make.


How would you make it, without contamination by NaClO, NaClO4, and NaCl?

blogfast25 - 19-6-2011 at 09:43

Quote: Originally posted by AJKOER  
My bias is an experimentalist as there are many reactions (especially with regard to catalyst) for which you would read that honest scientist admit that, as of yet, they do not understand the reaction path. Or, in your language, the current reaction is "THERMODYNAMICALLY UNFAVOURABLE", albeit, the reaction does somehow presumably occurs.


Yet the reactions you propose are all well understood to have large, positive ΔG. No catalyst (known or hidden) changes that: catalysts lower the kinetic barrier to a reaction but they don’t affect the ΔG of the reaction; if the reaction wasn’t already thermodynamically favourable (ΔG < 0) a catalyst can’t make it happen. See also Hess’ Law.
Quote: Originally posted by AJKOER  
Thanks for your opinion in largely discrediting Candidate I, which I, without calculations, was largely suspicious of, although I thought it was very interesting that a Ag/Al2O3 catalyst was able to produce a decomposition reaction on NH3 at 140 C in a gaseous state. As you may be aware of, ferrates can decompose aqueous NH3 at room temperature, but no reason to believe that they were at work here.


You ‘were largely suspicious of’, ‘without calculations’: do you know the heat of formation of Al2O3? Look it up. Oh go on: here it is: - 1,676 kJ/mol (off the top of my head). You think a bit of ammonia can break that down?

An Ag/Al2O3 catalyst is a very different thing than Al2O3 obtained either from burning or from aqueous precipitation.
Quote: Originally posted by AJKOER  
However, you did include the Candidate 2, in your no fly list. Do this mean you have a better path than I suggested?


No.
Quote: Originally posted by AJKOER  
Do you doubt the published research I cited (which came up under a search for Aluminum Nitride)?


The cited research is largely irrelevant to your central claim. What I think about the quoted researchers is immaterial.
Quote: Originally posted by AJKOER  
I performed this reaction many times and also place the reactants in a partially crushed plastic bottle, which under pressure from the gas generated, popped back into shape so be careful. One time the reaction was particularly interesting as the final reactant liquid was clear (excess NH4OH?) and I wanted to react the solution (NH4OH+Al(OH)3=?) with something else (it was MgSO4, very cool, forms a double precipitate Ammonium Aluminum Sulphate and Mg(OH)2 ), so I poured about half of the solution out leaving some still solid Al/Al2O3 in the bottle. The next day, the one time clear solution was cloudy with obvious jelly like Al(OH)3! I only poured some out! What I also found interesting is that some authors mention that excess NH4OH does produce a clear solution with Al(OH)3 and others, including yourself, say it doesn't. To quote: "This precipitate of Al(OH)3, which is amphoteric, dissolves in an excess of hydroxide or in acids". Source: http://www.public.asu.edu/~jpbirk/qual/qualanal/aluminum.htm...


Dear G-d, so many fallacies and spurious claims in one go!

NO, you haven’t ‘performed this reaction many times’. Nowhere have you proved what the gas you claim to have observed actually is. I suggest you take a clear carbonated drink, put some Al2O3 in it and shake it: do you believe the evolved gas is N2 + H2?!?!

“like Al(OH)3”: once again you provide no evidence whatsoever: it just ‘looks like Al(OH)3’. And still you call yourself an ‘experimentalist’?

To quote: "This precipitate of Al(OH)3, which is amphoteric, dissolves in an excess of hydroxide or in acids".

And where doest it say that Al(OH)3 dissolves in NH3 solution??? In your overheated imagination perhaps but not (of course not!) in the source text you link to. In fact it clearly implies that the precipitate (of Al(OH)3) obtained by adding NH3 solution to an aluminium salt can be dissolved only in an NaOH (or other strong hydroxide).

So why is this? Well, where you’ve been consistently writing an ammonia solution as “NH4OH” (sic), there really is no such thing. NH3 is very soluble in water but is a very weak base:

NH<sub>3</sub> (aq) + H<sub>2</sub>O (l) < === > NH<sub>4</sub><sup>+</sup> (aq) + OH<sup>-</sup> (aq)

… has a very small equilibrium constant (look it up; pK<sub>b</sub> of ammonia): the overwhelming majority of ammonia in solution is in fact present as NH<sub>3</sub>, NOT as NH<sub>4</sub><sup>+</sup> + OH<sup>-</sup>. This simple fact does not change much with higher concentrations of ammonia solution. That makes ammonia solution fundamentally different from NaOH (and similar) which does dissociate completely to Na+ and OH-.

To dissolve Al(OH)3 in alkali the (somewhat simplified) reaction path is:

Al(OH)3 (s) + OH<sup>-</sup> (aq) < === > Al(OH)<sub>4</sub><sup>-</sup> (aq)

For the reaction to proceed appreciably, the concentration of OH<sup>-</sup> (noted as [OH<sup>-</sup>]) needs to be sufficiently high. Even concentrated ammonia solutions do not reach the required [OH<sup>-</sup>] threshold. With NaOH on the other hand it is only a question of making the solution sufficiently concentrated.
Quote: Originally posted by AJKOER  
The source could be trustworthy as it is basically pictures of actual reactions. I personally am in the "it depends camp" meaning is there still Al or Al2O3 to be reacted with, pH, etc.

Thanks, however, for being skeptical, for this is how I believe knowledge progresses!


The source is trustworthy but it isn’t because of the pictures: for two different reactions they use the same photo!

m1tanker78 - 19-6-2011 at 10:19

Quote:
Even concentrated ammonia solutions do not reach the required [OH-] threshold.


Amen to that. I added some Al-laced Al2O3 left over from casting Al to 35% NH3 solution of high purity. The reaction proceeds just the same as using <5% household ammonia solution. The only real difference is that one makes you wrinkle your nose and the other punches you in the face!

There are a couple of conditions I'd like to alter before I call it quits.

Tank

blogfast25 - 19-6-2011 at 13:15

Quote: Originally posted by m1tanker78  
Amen to that. I added some Al-laced Al2O3 left over from casting Al to 35% NH3 solution of high purity. The reaction proceeds just the same as using <5% household ammonia solution. The only real difference is that one makes you wrinkle your nose and the other punches you in the face!

Tank


Which reaction precisely are you now referring to?

For a weak base in fairly dilute concentration, the hydoxide ion concentration is approx.: [OH<sup>-</sup>] = SQRT (K<sub>b</sub> . C<sub>b</sub>;)

with K<sub>b</sub> the Bronstedt base equilibrium constant and C<sub>b</sub> the molar concentration of the base (SQRT is square root). But for more concentrated solutions it plateaus off.

For strong bases,

[OH<sup>-</sup>] = C<sub>b</sub>

Edit:

The pK<sub>b</sub> = - log K<sub>b</sub> = 4.75 for ammonia.

A commercial household ammonia solution is about 4 % or just over 2 M (mol/L). The expected [OH<sup>-</sup>] ≈ √ (10<sup> -4.75</sup> x 2) = 0.006 or a pH of about 11.8.

For a similar NaOH solution (also 2 mol/L), [OH<sup>-</sup>] ≈ 2 and pH > 14!



[Edited on 20-6-2011 by blogfast25]

m1tanker78 - 20-6-2011 at 09:50

Quote:
Which reaction precisely are you now referring to?

To be honest, I don't know what the precise reaction is. I performed a simple experiment to qualitatively show what you quantitatively put forth.

Quote:
But for more concentrated solutions it plateaus off.

That was the name of the game. Your mathematical interpretation was very helpful in putting a name to the face - well, actually, a number.

It's somewhat difficult to test the evolved gas composition because there's so little of it. My guess would be H2 but with all that NH3 vapor around, even a liquid displacement gas measurement apparatus may give false results.

How do Al ions affect NH3 solubility in aqueous solution??

Tank

blogfast25 - 20-6-2011 at 12:53

Quote: Originally posted by m1tanker78  
How do Al ions affect NH3 solubility in aqueous solution??

Tank


Not at all: in water the NH3 protonates a little to NH4<sup>+</sup>, leaving the same amount of OH<sup>-</sup>. Any Al<sup>3+</sup> immediately reacts with the OH<sup>-</sup>, to hydrated alumina.

If you keep adding aluminium ions, it ends up like a displacement reaction:

Al2(SO4)3 (aq) + 6 NH3 (aq)+ 6 H2O (l) === > 2 Al (OH)3 (s) + 3 (NH4)2SO4 (aq)

LanthanumK - 20-6-2011 at 13:32

If anything, it increases its solubility as the equilibrium NH4+ + OH- <=> NH3 + H2O is moved to the left.

Neil - 20-6-2011 at 17:34

While watching Gert go off like a rocket is kinda fun...


I tried out this reaction as put forth using household ammonia with negative results. I tried heating the aluminum foil to several different oxide/metal levels and tested each successively and had no gas produced. So far this makes sense.


But I might have an olive branch,

If you add a bit of chloride(I used KCl) the solution does produce a very fine bubbled gas.

I heated the test tube in a water bath and while the gas production did increase it was was still slight.


So, in my experiments I could not produce any sort of gas with ammonia and aluminum foil no matter how well it was 'burnt'.

I added a chloride source (KCl) and gas was produced.

@ m1tanker78; Do you use a chloride for flux?

AJKOER - 20-6-2011 at 18:36

These comments have promoted some research on my part that many may find interesting. One article by Prideaux and Henness 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."
SOURCE:
http://pubs.rsc.org/en/content/articlelanding/1940/an/an9406...

My prior understanding on complex ion behavior is that, for the most part, it can be explained based on Lewis acid base theory. For example, in the reaction:

NH3 + Cu (2+) --> Cu(NH3)4 (2+)

the NH3 with an unshared electron pair, is electron rich and can donate an electron pair (act as a Lewis base), whereas, the Cu+2 is electron deficient and can accept an electron pairs (act as a Lewis acid). Some key conditionals for complex ion formation being "concentrated" and "excess".

On the ionization of NH4OH discussion, the influence of a chloride ion and the existence of ammonim aluminate, here is an excellent discussion in "Journal of the American Chemical Society", Volume 38, page 1287 (albeit a bit dated):

"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 A1(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&...

[Edited on 21-6-2011 by AJKOER]

blogfast25 - 21-6-2011 at 05:27

1916, fantastic vintage! Next we’ll be reviving Phlogiston Theory to try and explain the basically non-existent ammonium aluminate. :D
Quote: Originally posted by Neil  
If you add a bit of chloride(I used KCl) the solution does produce a very fine bubbled gas.

I heated the test tube in a water bath and while the gas production did increase it was was still slight.


Try and make a bit more and see what it is. My money is firmly on good ole’ NH3…




[Edited on 21-6-2011 by blogfast25]

Neil - 21-6-2011 at 05:50

Quote: Originally posted by blogfast25  
1916, fantastic vintage! Next we’ll be reviving Phlogiston Theory to try and explain the basically non-existent ammonium aluminate. :D

Quote: Originally posted by Neil  
If you add a bit of chloride(I used KCl) the solution does produce a very fine bubbled gas.

I heated the test tube in a water bath and while the gas production did increase it was was still slight.


Try and make a bit more and see what it is. My money is firmly on good ole’ NH3…




[Edited on 21-6-2011 by blogfast25]




My money is that you're at least 90% right. The ammonia solution of last night has almost no ammonia in it this morning. I do think the remaining Al in the "burnt foil" was reacting and producing hydrogen as Al does in Ammonium Chloride solutions, but I suppose I'll have to test it won't I?


More to the point; Aluminum is used to store and handle ammonia hydroxide, ammonia gas, ammonia sulfate and ammonia nitrate - at least according to the engineering tables that come up if you Google 'ammonia corrosion table'

If ammonia reacted with Al2O3 we would be using it to dissolve slag and concentrate bauxite... We sure wouldn't be using it to store ammonia.

Phlogiston would have made for a much more interesting world...

AJKOER - 21-6-2011 at 06:12

Actually, I find fascinating the tenacity of theorists!

I agree with the safety of Al container storage for NH3. However, this is not the case for NH3+H2O (I read on a NASA site once of a NH3/Al tubing system that also mentioned in the event of moisture entering the system, a corrosive reaction could occur).

The sad truth is that anyone who ever has owned anything Aluminum knows (and a 2 minute internet search more than confirms) you cannot clean Aluminum products with household ammonia (it causes pits).

As a chemist, you should know that this means that household ammonia (or its detergents? really) therefore dissolves/attacks the protective Al2O3 coating (or the 2% impurities which vary from Si, Fe, Mg and Mn?). This most likely means a chemical reaction given the resistance of protective layer. One model (the least complex given the near entirely with which Al eventually dissolves) is the formation of an Aluminum ammonium complex that exposes the underlying Al which readily reacts with H2O to produce H2 (Source: See "Concise Encyclopedia Chemistry" by deGruyter or search the web). An undisclosed source on another forum states that the NH3 acts as a catalyst in the dissolving of Al with NH4OH with the release of H2 (which was confirmed by the forum's author). I will see if I can obtain the actual source, but I am not expecting more than a casual reference.

I would advise you read the google book's entire section on Al2O3 and the other article that dates from 1940, good stuff!

[Edited on 21-6-2011 by AJKOER]

m1tanker78 - 21-6-2011 at 07:04

Neil, to answer your question, I didn't use chloride flux in the melt. In fact, I didn't add anything at all except Al bits of high purity. These high purity ingots will be used to make 'designer' alloys at a later date so I accepted a slight loss of Al in the slag.

That's also the reason I chose this lot for this experiment. I wanted to remove as many ?'s as possible. The surfactants they add to household ammonia drive me up the wall (almost as much as the smell). Also, Al foil is far from pure and has tricked me in the past.

I'm trying to digest everything since Al chemistry is mostly new to me. I have a clear solution that has turned somewhat viscous (not gelatinous). If I filter the bits out and heat the solution, NH3 will escape along with water. Eventually, Al(OH)3 will lose water, first forming AlO(OH), and finally, Al2O3. If that's correct then why can't this be exploited as a path to quantifying the amount of Al present in slag or in AJOKER's case, burnt aluminum? I don't believe Al2O3 reacts appreciably (if at all) with NH4OH or NH3 in aqueous solution. I know that stronger alkalis could be used but the attractive side of ammonia is that it completely evaporates, allowing a sample to be dried and weighed without significant influence from other ionic salts.

One thing that's bugged me for a while...

Wiki's Al2O3 page says alumina is insoluble in water - OK. It also says Al2O3 is 'very hygroscopic'. Is this a typo? I've always been under the impression that alumina is extremely stable in air and moisture.

Tank

Neil - 21-6-2011 at 08:48

Quote: Originally posted by AJKOER  
Actually, I find fascinating the tenacity of theorists!

I agree with the safety of Al container storage for NH3. However, this is not the case for NH3+H2O (I read on a NASA site once of a NH3/Al tubing system that also mentioned in the event of moisture entering the system, a corrosive reaction could occur).

The sad truth is that anyone who ever has owned anything Aluminum knows (and a 2 minute internet search more than confirms) you cannot clean Aluminum products with household ammonia (it causes pits).

As a chemist, you should know that this means that household ammonia (or its detergents? really) therefore dissolves/attacks the protective Al2O3 coating (or the 2% impurities which vary from Si, Fe, Mg and Mn?). This most likely means a chemical reaction given the resistance of protective layer. One model (the least complex given the near entirely with which Al eventually dissolves) is the formation of an Aluminum ammonium complex that exposes the underlying Al which readily reacts with H2O to produce H2 (Source: See "Concise Encyclopedia Chemistry" by deGruyter or search the web). An undisclosed source on another forum states that the NH3 acts as a catalyst in the dissolving of Al with NH4OH with the release of H2 (which was confirmed by the forum's author). I will see if I can obtain the actual source, but I am not expecting more than a casual reference.

I would advise you read the google book's entire section on Al2O3 and the other article that dates from 1940, good stuff!

[Edited on 21-6-2011 by AJKOER]


Your condescending attitude is very much appreciated, thank you.

Household ammonia should not be used to clean aluminum ergo Ammonia causes pitting to Aluminum by removing its oxide coat is absurd. That is a logical fallacy based on far to little knowledge.

Cleaning implies you are removing a unwanted substance, what are the chances that 100% of the time that an aluminum surface, which is in need of cleaning, will have no traces of halites or phosphates on it?

What are the chances that relatively crude cleaning ammonia is free of any contaminates or do you clean your counters with analytical grade ammonia?

To be trite, I will now preform an experiment following your methodology. I saw a man who resembled Elvis, to verify that this is Elvis I will now Google "Elvis is alive, proof, pictures" I now have have proof Elvis is alive.

The undisclosed forum to which you refer, and to which you post the entries you post here verbatim, does not by mystical fortune confer objectivity or legitimacy to your claims in and of its self.


As a student of the Humanities and of Science I know that anything can appear true if you solely look for validation.

As a Student of Chemistry I know that chemists are individuals who have well studied and examined many reactions in many situations. I know from experience that Gert is exceptionally good with his mathematical analytical skills. I also know that he is not right 100% of the time and that he does not know everything, who does? While he may deliver his opinions with a sledge hammer, as a student of chemistry I know that they carry far more weight then my own deductions by virtue of his experience and education.


I refer you to this

http://books.google.ca/books?id=0CYdprLzfMoC&pg=PA611&am...

This
http://www.usmotors.com/Products/ProFacts/tableof.htm

This
http://www.springerlink.com/content/jj5t8210k44g1u33/

This
http://www.dynonobel.com/files/2010/04/1130-Aqua-Ammonia-09-... (what aluminum is not an incompatible?)

This
http://home.comcast.net/~pchristou/docs/silhoutes.pdf

and so on.


Here is a test, take aluminum foil and blow torch it with an oxidizing butane flame and then test for residual Al with NaOH.

No matter how much fuel you waste you will never get all of the aluminum to convert to an oxide. The Aluminum has to high of a boiling point to fume out and burn, and the alumina has to high of a melting point to let the aluminum out to oxidize. if you use Draino, the reaction with Al will produce some ammonia gas which you can test for with litmus paper.

So, lets say that Al2O3 may have nothing to do with the observed reaction. What happens if we rough up some Al foil and toss it in Ammonia and heat the ammonia? Nothing. What about if we take the sand paper and rub off some of the garnet sand and add that? Nothing. so aluminum and ammonia did not react, nor did garnet and ammonia. Ladies fear not, Windex will not dissolve your rubies.

What if we add some KCl? slow gas production is observed originating from the scratches on the Al foil. increased heat causes an increase in reactions.

Now, I did test the gas from the burnt Al foil. I washed it with HCl, which did not become cloudy and did not dissolve any of the gas. A flame test gave a very energetic fire ball, which suggests Hydrogen.

I got the exact same reaction for plain aluminum foil in the same ammonia solution with KCl.


So where does that leave us?

http://proj3.sinica.edu.tw/~chem/servxx6/files/paper_7425_12...
From above:
"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."

There are many routes to pitting Al that do not involve digesting its oxide. Without making an effort to determine if any of these are active the assumption that the ammonium is somehow reacting with the oxide is baseless and poor science.

What is interesting is that using pure ingredients m1tanker78 has what is at least a very similar reaction. (One last grasp for a straw, did you use a de-gassing agent?)

I did not observe any cloudy formations when the captured gas was brought near HCl, which suggests that it was ether consumed or otherwise diverted from the captured gas.

What seems more likely to me is that when Al is heated to generate a thick oxide skin, the skin is more vulnerable to cracking which exposes fresh metal. The fresh metal would be able to react with the water producing hydrogen and possibly a complex as it transitions to becoming an insoluble hydroxide which could form a gel.

As to Al2O3 my understanding is that it's hygroscopic tendencies disappear once it has been calcined similar to the tendency of CaSO4 to cease absorbing water once it has been calcined at more then 450F. Both substances experience crystal phase changes which create a more water resistant form.

m1tanker78 - 21-6-2011 at 09:59

Quote:
What seems more likely to me is that when Al is heated to generate a thick oxide skin, the skin is more vulnerable to cracking which exposes fresh metal. The fresh metal would be able to react with the water producing hydrogen and possibly a complex as it transitions to becoming an insoluble hydroxide which could form a gel.

This could certainly be the case. However, my belief is that the solvent (aq. ammonia) is simply migrating through the porous oxide - very slowly. After all, heavy Al anodizing is performed in this way but usually in acidic solution. Regardless, the liquid must migrate through the pores somehow. This probably also contributes to the observation of near equal rates of reaction in dilute ammonia water and concentrated ammonia water.

Neil, it's worth mentioning that Al foil manufacturers usually coat the foil in a protective wax (or other similar substance). Other than that, I don't have a good explanation of why your foil didn't produce any bubbles. My observations are identical to what AJOKER described in the OP. Try tearing the foil into small pieces and look for bubbling at the edges.

Tank

AJKOER - 21-6-2011 at 11:18

I found an interesting reference on one of the slides titled "Presence of Aluminum Nitride in Salt Cake" presented by the Global Symposium on Recycling, Waste Treatment and Clean Technology in Ocober 2008, Cancun, Mexico. In essence, the dissolving of AlN produces ammonia that raises the pH which "dissolves the alumina film on unrecovered Aluminum particles surface, thereby exposing the Al surface to the reaction Al + H2O --> Al(OH)3 + H2"

One could read this as the NH3 is directly responsible for dissolving the Al2O3 as the pH is raised, or as the salt cake has NaCl, KCl, MgAl2O4 and oxides of Fe, Si and Zn, these in the presence of a higher pH are responsible.

Link: http://www.es.anl.gov/Energy_systems/docs/process_tech/indus...

As a sidebar, I also observed that leaving Al foil in vinegar for a few hours, even if apparently unreacted, does speed up the speed of the reaction after the Al foil has been wash and placed into NH4OH. I got this trick from an Al Foil Coating manufacture article that lamented the power of acetic acid to pierce through all their efforts to make the Al foil unreactive.

Here is the link on the NASA use of anhydrous NH3 with Al piping:
http://www.nasa.gov/offices/oce/llis/0698.html

Also from the cited reference: "Corrosion resistant materials handbook By D. J. De Renzo, Ibert Mellan", page 611, it appears with respect to the reaction of NH4OH on Al Alloys that "ammonium hydroxide has a rapid initial reaction on aluminum alloys which decreases dramatically as pH and concentration increase". Some reasons given relating to Al film formation.

For those rejecting a direct NH4OH and Al2O3 reaction, a conjecture is a possible reaction path with NH3 facilitating the creation of say, FeAl2O3 (or other impurity). See "Alumina as a textural promoter of iron synthetic ammonia catalysts" by H. Topsøe, J. A. Dumesic and M. Boudart. However, reaction with Aluminum free of Fe impurity would nullify this hypothesis (which, by the way, I rate as unlikely).


[Edited on 21-6-2011 by AJKOER]

[Edited on 22-6-2011 by AJKOER]

Neil - 21-6-2011 at 12:19

Quote: Originally posted by m1tanker78  

Neil, it's worth mentioning that Al foil manufacturers usually coat the foil in a protective wax (or other similar substance). Other than that, I don't have a good explanation of why your foil didn't produce any bubbles. My observations are identical to what AJOKER described in the OP. Try tearing the foil into small pieces and look for bubbling at the edges.

Tank


The wax is a very good point but I didn't observe any reaction with the "burnt" aluminum or the roughed aluminum until I added the salt. It looks like the fractures in the "burnt" foil seem to produce the most bubbles.



blogfast25 - 21-6-2011 at 12:59

Quote: Originally posted by AJKOER  
Actually, I find fascinating the tenacity of theorists!

I agree with the safety of Al container storage for NH3. However, this is not the case for NH3+H2O (I read on a NASA site once of a NH3/Al tubing system that also mentioned in the event of moisture entering the system, a corrosive reaction could occur).

The sad truth is that anyone who ever has owned anything Aluminum knows (and a 2 minute internet search more than confirms) you cannot clean Aluminum products with household ammonia (it causes pits).

As a chemist, you should know that this means that household ammonia (or its detergents? really) therefore dissolves/attacks the protective Al2O3 coating (or the 2% impurities which vary from Si, Fe, Mg and Mn?). This most likely means a chemical reaction given the resistance of protective layer. One model (the least complex given the near entirely with which Al eventually dissolves) is the formation of an Aluminum ammonium complex that exposes the underlying Al which readily reacts with H2O to produce H2 (Source: See "Concise Encyclopedia Chemistry" by deGruyter or search the web). An undisclosed source on another forum states that the NH3 acts as a catalyst in the dissolving of Al with NH4OH with the release of H2 (which was confirmed by the forum's author). I will see if I can obtain the actual source, but I am not expecting more than a casual reference.

I would advise you read the google book's entire section on Al2O3 and the other article that dates from 1940, good stuff!

[Edited on 21-6-2011 by AJKOER]


AJKOER:

When in a hole stop digging.

You call me a ‘theorist’, yet you the ‘experimentalist’ have yet to provide the first shred of evidence for your wild theories, never mind the ill digested sources which you keep rolling out, the latest one: ‘An undisclosed source on another forum states that the NH3 acts as a catalyst in the dissolving of Al with NH4OH with the release of H2 (which was confirmed by the forum's author). I will see if I can obtain the actual source, but I am not expecting more than a casual reference’.

This is also priceless: “One model (the least complex given the near entirely with which Al eventually dissolves) is the formation of an Aluminum ammonium complex”.

So someone who is unable to prove what the observed gas really is now comes up with another corker: an ‘aluminium ammonia complex’. There is no such thing but that never stops you, does it?

I could go on but I’ll limit myself to: ‘and what Neil said…’

You remind me of those ‘inventors’ of perpetual motion machines who say: “yeah, I know about the Second Law of Thermodynamics but I feel my latest design will probably work anyway’.

[Edited on 21-6-2011 by blogfast25]

AJKOER - 21-6-2011 at 14:15

Hi:

Please see my revised note above that could support my argument that NH3 is dissolving the Al2O3 (although, to be fair, it is the added NH3 from the AlN that is alluded to), or it is the chlorides and other components in the Salt Cake, in a higher pH environment, that are responsible. Interesting, these are the two hypothesis currently on the table (a direct NH3 on Al2O3 path, or a corrosive Chloride activated reaction on the Al2O3, both in the presence of an elevated pH.

I would like to amend my statement that the dissolving of Al2O3 in the presence of NH3 (one possible path), does not necessarily means the creation of NH4[Al(OH)4], it may be that other unknown mechanisms are at work. I can also accept the Chloride activation path as the cause of the weakening of the Al2O3 coat.

[Edited on 21-6-2011 by AJKOER]

[Edited on 22-6-2011 by AJKOER]

Neil - 21-6-2011 at 17:53

Quote: Originally posted by AJKOER  
I found an interesting reference on one of the slides titled "Presence of Aluminum Nitride in Salt Cake" presented by the Global Symposium on Recycling, Waste Treatment and Clean Technology in Ocober 2008, Cancun, Mexico. In essence, the dissolving of AlN produces ammonia that raises the pH which "dissolves the alumina film on unrecovered Aluminum particles surface, thereby exposing the Al surface to the reaction Al + H2O --> Al(OH)3 + H2"

One could read this as the NH3 is reputedly directly dissolving the Al2O3 as the pH is raised.

Link: http://www.es.anl.gov/Energy_systems/docs/process_tech/indus...

As a sidebar, I also observed that leaving Al foil in vinegar for a few hours, even if apparently unreacted, does speed up the speed of the reaction after the Al foil has been wash and placed into NH4OH. I got this trick from an Al Foil Coating manufacture article that lamented the power of acetic acid to pierce through all their efforts to make the Al foil unreactive.

Here is the link on the NASA use of anhydrous NH3 with Al piping:
http://www.nasa.gov/offices/oce/llis/0698.html

Also from the cited reference: "Corrosion resistant materials handbook By D. J. De Renzo, Ibert Mellan", page 611, it appears with respect to the reaction of NH4OH on Al Alloys that "ammonium hydroxide has a rapid initial reaction on aluminum alloys which decreases dramatically as pH and concentration increase". Some reasons given relating to Al film formation.

For those rejecting a direct NH4OH and Al2O3 reaction, a conjecture is a possible reaction path with NH3 facilitating the creation of say, FeAl2O3 (or other impurity). See "Alumina as a textural promoter of iron synthetic ammonia catalysts" by H. Topsøe, J. A. Dumesic and M. Boudart. However, reaction with Aluminum free of Fe impurity would nullify this hypothesis (which, by the way, I rate as unlikely).


[Edited on 21-6-2011 by AJKOER]




:(
Quote:

In essence, the dissolving of AlN produces ammonia that raises the pH which "dissolves the alumina film on unrecovered Aluminum particles surface, thereby exposing the Al surface to the reaction Al + H2O --> Al(OH)3 + H2" One could read this as the NH3 is reputedly directly dissolving the Al2O3 as the pH is raised.




YES! EXACTLY! One could read it as that, if one wants to imagine chemical reactions, completely ignore the scientific process, misquote and misrepresent a source and generally make up information to satisfy a theory which has no basis.

What it actually says is that the production of ammonia contributes to a rise in pH which aids in the digestion of aluminum still trapped in the salt cake which was extracted from aluminum slag which is riddled with Chlorides.....

Here is the quote


Quote:

"Presence of aluminum nitride in salt cake
•
Reactivity with water
–
AlN + H2O Al2O3 + NH3
–
NH3 + H2O NH4OH (pH increases)
–
High pH dissolves alumina film on un-recovered aluminum particle surface
•
Exposes aluminum surface to reaction
–
Al + H2O Al(OH)3 + H2 (+ heat)
–
Hot H2 + O2 (air) + combustibles fire"



He is specifically talking about material which he described on the previous page:
Quote:

"Why recycle salt cake?
•
Recover residual aluminum (4-8%)
–
Economic and environmental benefit
–
Recover more that half of residual Al
•
>50% Al energy content (break-even at 2-3% Al)
•
Perceived environmental hazard related to salt cake composition
–
Reactive, pyrophoric
–
Noxious ammonia odor
–Leachable chloride content
"


Bold added for emphasis.

:o if your household ammonia sounds like the above mixture you may want to call HazMAt and get a lawyer. Chlorides are long established as eating Al for breakfast - nothing in that presentation backs up anything you've claimed.


Suggesting that the author of the paper, who it seems is a noted expert on aluminum and magnesium chemistry, is saying that ammonia dissolves alumina when he suggests no such thing is rather... bad form
Quote: Originally posted by AJKOER  
Also from the cited reference: "Corrosion resistant materials handbook By D. J. De Renzo, Ibert Mellan", page 611, it appears with respect to the reaction of NH4OH on Al Alloys that "ammonium hydroxide has a rapid initial reaction on aluminum alloys which decreases dramatically as pH and concentration increase". Some reasons given relating to Al film formation.



"decreases dramatically as pH and concentration increase" :o


From the NASA page you linked on heat exchangers, under a graph in which they delineate water and aluminum as being incompatible for the purposes of aluminum heat exchangers is this sole mention:
"trace amounts of water in ammonia can lead to a reaction with the aluminum container and the formation of hydrogen gas."


Hey what is this? http://www.thescienceforum.com/viewtopic.php?t=16168&pos...

Relivance? You consider someone, who thinks dipping Al metal into "90% H2O, 10% NH3" creates AlN, as a credentialed source?


And this?
http://www.sciencemadness.org/talk/viewthread.php?action=pri...
Quote:
AJKOER here yet again. My personal observation is that freshly burned Al foil is more reactive than the original foil with respect to action by NH4OH.



So burnt aluminum reacts better then unburnt aluminum?

Nope wait a sec


Quote:
Also, only the edges of torn foil are producing bubbles with NH4OH and after time, not all the foil is dissolving (some foil strips remain resistance and are untouched).



So freshly exposed aluminum which would likely have little oxide cover reacts faster, therefore alumina is dissolved by ammonia?

You change your story, forget data as it becomes evident that it contradicts you and ether intentionally or unintentionally misrepresent sources and make spurious claims which are not only impossible but are down right ignorant and then you have the gall to patronize and insult others because you are unwilling to admit you may be mistaken?

You've dug so deep you hit bull shit. You may find this helpful
http://www.envirochem.co.nz/pdfs/cowmate_techdata.pdf


Still interesting are the results m1tanker78 produced. The only aluminum I have access to is all alloyed, I'd like to know more and am trying to find a source of low alloy Al to repeat your reaction.

AJKOER - 21-6-2011 at 18:00

Yes, upon re-reading the whole slide presentation, I independently revised my statement, because I wanted to be fair in the event someone else did not pay attention to the other slides.

Sorry for the delay. I am basically in agreement with a two hypothesis scenario.

One of my problems is rejecting the observation that two different household ammonia actually do dissolve both Al and Al2O3. Neither of these NH4OH solutions leave any salt residues! Please note that if you have read the tread, at least one other person has confirmed some of my observations and at least one person has mentioned that the NH3 may be invading the porous Al2O3.

By the way, I noted that vinegar treated Al foil and Al2O3 both dissolved in NH4OH at about the same rate. This makes sense in that burning the Al removes plastic coatings, oils and the annealing of the Al, all of which increases Al resistance to acids and the like.



[Edited on 22-6-2011 by AJKOER]

Neil - 21-6-2011 at 18:03

You can do this at home;

http://apps.caes.uga.edu/sbof/main/lessonPlan/BasicSynthesis...

This would also serve you well:

http://library.thinkquest.org/2923/tests.html

AJKOER - 21-6-2011 at 20:27

I think we close to a resolution.

First, here is an important reaction (I have performed it myself) that used in the mining of Copper ore:

CuO + 4 NH4OH -- Ammonium carbonate---> Cu(OH)2.(NH3)4 + 3 H2O

That is, the CuO is commercially successfully dissolved by aqueous ammonia in the presence of a weak acid or ammonium carbonate. Interestingly, Concise Encyclopedia Chemistry by DeGruyter notes that (NH4)2CO3 solution dissolves CuO.

Second, assume we have a general consensus that Al2O3 and NH4OH dissolve in the presence of a NaCl or KCl (we have a reference and direct observations on this).

So per our Copper ore example, what compound would you think is produced from the reaction of Al2O3 in excess NH4OH in the presence of NaCl?





[Edited on 22-6-2011 by AJKOER]

Fleaker - 21-6-2011 at 20:41

What reaction?


Al2O3 has different modifications. Some are easier to digest than others. Ammonia won't do the job well at all.

AJKOER - 21-6-2011 at 21:52

The Al203/AlN/Al prepared from burning Al with traces of Si and Fe, for example.

blogfast25 - 22-6-2011 at 05:15

Neil:

Excellent debunking but you won’t really catch AJKOER: he moves the goal posts al the time. Take this for example:

“Please note that if you have read the tread, at least one other person has confirmed some of my observations and at least one person has mentioned that the NH3 may be invading the porous Al2O3.”

Regards the latter part of that statement, it was Tank who correctly pointed to diffusion of NH3 through the Al2O3 based passivation layer. That is indeed very likely to be the case. Take titanium for example: like Al it passivates to a TiO2 layer of only a few atoms thick. TiO2 is COMPLETELY insoluble in HCl, no matter how concentrated or hot (with the exception of very freshly prepared Ti(OH)4), yet hot, strong HCl attacks titanium metal readily, so vigorously that the acid is used to make Ti<sup>3+</sup> solutions (I’ve done this many times). Obviously the acid (like NH3 a small molecule) can penetrate the oxide layer w/o actually dissolving it.

Yet the partial acceptance of Tank’s claim is now presented as if it confirms the reaction between “NH4OH” (sic) and alumina.

With regards to the dissolution of Al2O3, I suggest AJKOER reacts some Al powder with KClO3: 2 Al + KClO3 === > Al2O3 + KCl. This runs so hot you obtain annealed alumina (the KCl get blown off because the temperature exceeds its boiling point). I suggest he then tries to dissolve the alumina in the alkali of his choice: whether “NH4OH”, NaOH or KOH. Alternatively look up ‘Bayer process dissolution Bauxite’ to grasp just how difficult the digestion of alumina with strong alkalis really is. It might just about shut him up about “NH4OH” in this context.

AJKOER - 22-6-2011 at 05:50

Thanks for the interesting points blogfast25, but the question does not relate to the dissolving of Alumina by NH4OH at this time, but Alumina + Salt + NH4OH, with the clear understanding that it is the burnt Al contaminated with AlN and Al (that is, the one we have been experimenting with) in play.

Also, there is no time limit on this game (forgive the sport analogy, but blogfast25 started it), so a boring game (you may have seen one) with a point on the scoreboard taking forever (meaning that no one would be likely to sell this game commercially) still counts.

Still waiting for an answer.

blogfast25 - 22-6-2011 at 06:03

Quote: Originally posted by AJKOER  
Still waiting for an answer.


To what question? You’ve changed the story that many times hardly anyone here still knows what you’re talking about! It’s certainly been quite an arc: from alumina cracking NH3 into nitrogen and hydrogen to mixtures of Al2O3 and AlN giving rise to NH3 and the latter assisting in the dissolution of residual Al.

So when, instead of trying to unsuccessfully dodge every bullet, will you present some actual observations, followed up perhaps with some real evidence for what you observed, huh? How about that?

m1tanker78 - 22-6-2011 at 06:26

Quote: Originally posted by blogfast25  
Quote: Originally posted by AJKOER  
Still waiting for an answer.
[...]present some actual observations, followed up perhaps with some real evidence for what you observed, huh? How about that?


I agree. AJOKER, what is it you're after? I second Blog's suggestion to look into the Bayer Process. Set up some meaningful experiments and get some rewarding 'hands-on' time. As you can see, it's a tough crowd! :D I, myself, learned that the hard way with some sloppy experiments and bad assumptions.

Tank

LanthanumK - 22-6-2011 at 06:30

Aluminium is still a metal; although its oxide is amphoteric it does not have the acidic character of a nonmetal oxide such as sulfur trioxide. Ammonia is a weak base, not a strong base. There is a difference between a strong and a weak base. Ammonia is not basic enough to dissolve Al2O3.

Copper oxide forms an ammine complex with NH3 ions. I don't know if this has been brought up before, but here is a thread about aluminium ammine ions: http://www.sciencemadness.org/talk/viewthread.php?tid=10977. Is this what you mean when you are referring to aluminium ammonia complexes?

PHILOU Zrealone - 22-6-2011 at 06:34

Maybe this observation will help?

Al does dissolve into NH4NO3.

Normally Al is said to be resistant to concentrated HNO3 and I wanted to test if I could make some Al(NO3)3 or a Al(NO3)3.xNH3 complex from Al and NH4NO3 solution... so I have put some Al foil into a glass jarr with a plastic screewcap and some concentrated NH4NO3 water solution.

At normal ambiant temperature nothing seemed to happen for days...
But about a few weeks later I thought back to that forgotten experiment bottle...and all the silvery Aluminium metal had dissappeared into a fine white mud...the cap was deformed and had tiny cracks on the side...I unscreewed the cap to allow somme gas to come out.

To my understanding ammonium/ammonia salts eather:
1°)Free some ammonia that allow partial or total dissolution of the Aluminium oxyd/hydroxyd layer in a reversible fashion.
2°)Make a soluble complex of the type NH4Al(OH)3(NO3)
3°)Make a transient (NH4)3AlO3 aluminate

The process seems slow but naked Al is produced because of gas generation and silvery foil dissappearing.
Al + 3H2O --> Al(OH)3 + 3/2 H2

I'm quite sure that NH4Cl will do even better (see my answer in the other tread on Cu(2+) and Al).

Reaction of NH4Cl + CuSO4 + Al must be even faster than with NaCl + CuSO4 + Al.
To add to the joy of the mix NH3 will complexate some Cu(2+) and CuO protective layer
and dissolve (with help of oxygen from the air) the just precipitated Cu powder.

AJKOER - 22-6-2011 at 07:18

With respect to "sloppy" experiments, I am not alluding to any of my work here, but to others in this thread using pure ingredients, better lab equipment and undoubtedly more seasoned lab procedures (thanks for the educational threads, I am sure many may benefit from them). Also, it was not also my suggested route of using KCl as a reaction catalyst for Alumina and NH4OH, so any criticism of "my" reaction or "my" sloppiness are not relevant.

The question has already been very clearly presented. If you have issues with the possible answer, or just don't know, please do not respond. I would respectfully ask you to display courtesy in allowing new or existing participants to contribute.

After receiving replies, I will revise my current views, or may just present the consensus view, and post to allow for final comments. Note, non-responses (including requests for time-outs, instant replays and/or attacks on the referee) will not contribute to the consensus view (or modify mine).

I sincerely wish to thank those who expended their time and lab resources to add to this tread.

AJKOER - 22-6-2011 at 07:43

PH Z (PHILOU Zrealone)

Thanks for prior experimental results that interestingly mirror my cited CuO + NH4OH which also proceeds in the presence of dilute acids (think HNO3) as I noted.

In the current case, think of Al2O3 + NH4OH ----in dilute HNO3-->

I like all of your answers, and I believed (despite some claims of my sloppiness), have witness the reversibility phenomena.

Thanks again.

Neil - 22-6-2011 at 08:10

Quote: Originally posted by AJKOER  
With respect to "sloppy" experiments, I am not alluding to any of my work here, but to others in this thread using pure ingredients, better lab equipment and undoubtedly more seasoned lab procedures (thanks for the educational threads, I am sure many may benefit from them). Also, it was not also my suggested route of using KCl as a reaction catalyst for Alumina and NH4OH, so any criticism of "my" reaction or "my" sloppiness are not relevant.

The question has already been very clearly presented. If you have issues with the possible answer, or just don't know, please do not respond. I would respectfully ask you to display courtesy in allowing new or existing participants to contribute.

After receiving replies, I will revise my current views, or may just present the consensus view, and post to allow for final comments. Note, non-responses (including requests for time-outs, instant replays and/or attacks on the referee) will not contribute to the consensus view (or modify mine).

I sincerely wish to thank those who expended their time and lab resources to add to this tread.



Thank you Gert.

AJKOER
Please do not suggest that I used KCl in an attempt to dissolve Alumina, I did no such thing. The problem is not that you are theorizing or attempting to further your understanding it is that you are making fallacious claims.

There are no scores in science, because ultimately there is only reality. Chemistry is nothing more or less then the study of the material which makes up your reality, matter. Scientist who decide there is a score cease being scientists in the same moment, for example Dr. Hwang Woo-Suk.

Contribution implies offering information, quid pro quo, ideally.


Do you have HCl, the ability to get red cabbage, Al foil, NaCl, household ammonia and distilled water?

What glassware do you have access to? Do you have a scale? Do you want a answer that you can verify on your own?

m1tanker78 - 22-6-2011 at 08:33

ajoker: get it together! Go back and read my post a little more carefully. Now, tell me who I accused of performing sloppy experiments. I'll be damned, I incriminated... myself.

For the record (and without double meanings, score keeping or any BS), I will say that you should perform and VERIFY some of YOUR OWN experiments when a theory flies into your head or when you read something that interests you. Feedback is good but don't expect everyone to drop what they're doing so they can pull the weight...

Tank

AJKOER - 22-6-2011 at 09:42

Neil:

Points well taken.

However, should I refer to the use of a catalyst (KCl) in the reaction of NH4OH on Alumina as your "reaction"?

Or, joint ownership? Or, product of thread participants work?

Or if you wish, I will, by default, assume full responsibility for it as I am aware of the Ammonium carbonate catalyst in the dissolving copper ore, and my dated reference does note the significance of dilute NH4Cl on Al(OH)3 precipitation, so I am comfortable with the concept.

Your choice, I would not want to ascribe any research to you without your consent if it was unintentional on your part.

Neil - 22-6-2011 at 10:24

"What color is the house on that hill?"






"Yellow!"


No!


"It appears yellow on this side!"

Yes!


Do you see the difference? In one instance I make an assumption, in the other I make an observation.

I observed nothing which suggests a reaction along the lines you are attributing.

I knew from experience that CuSO4 and Al react in water but need a Cl- source, normally I use NaCl.

Based on this knowledge I added a Cl- source. Why did I use KCl? Because I had no NaCl at hand.

What did I observe? A reaction occurred which produces a gas.


So to re-phrase:

The reason I added the Cl- source was because I hypothesized that it would allow the aluminum to react in a manner similar to the observations you described and that Cl- contamination in household ammonia was a not unlikely possibility.

What did I learn and observe?

The gas tested positive as being at least partially hydrogen. It tested negative for being ammonia. To be very fair, it tested positive for being a buoyant flammable gas. The KCl did in fact seem to trigger a reaction.

Did the KCl trigger the reaction? I do not know that for certain. The KCl came from NoSalt which contains anti caking agents which include tartrates and silicates IIRC. Could they have been responsible? Strictly speaking - I do not know.

If you wish to form a hypothesis that the observed reaction involved the Al2O3, which we assume to be present, then you need to find a way to measure any changes to the Al2O3 skin.

What about AlN? Do we know there is any present? Do we know if it is statistically significant? Do we know if it's formation is even plausible?

There are two sides to the scientific method. One is that it produces research which can be referred to and the other is that it can keep you alive.

For example, the post by AquaRegia where he detailed a very devastating event which occurred because he made a single assumption. You can never assume. In chemistry assuming doesn't just make an ass out of U it can also make you bleed or die.

I bring this up because based on other threads where you have posted it seems you have an interest in energetics. Based on what you've said and done in this thread, I strongly suggest you put in some textbook time. If you want to learn chemistry you have to be willing to trash a theory when it is undercut and you have to be open to being 100% wrong, otherwise you will end up being ether 100% ignored or worse, 100% dead.


If you wish to stop assuming, then start testing out what you suspect in a safe and scientific manner. If not, I suggest you take up Magic cards.

blogfast25 - 22-6-2011 at 11:50

Nice to see AJKOER has a crackpot supporter, one of the type that thinks if you can balance it, it must be chemistry:

”2°)Make a soluble complex of the type NH4Al(OH)3(NO3)
3°)Make a transient (NH4)3AlO3 aluminate”

O-kaaayy…

I’ve got a balanced equation for you:

(NO6)2U3BPZr5 + 3 H9SCl4O3 === > (NO3)2 + 3UZr + 2 ZrO2 + 12 HCl + P3O17 + 11/2 O2 + 3 H2 + BS3

Did I say BS?

AJKOER - 22-6-2011 at 12:21

Neil thanks for the picture.

Talk about trashing theories, I think we can start with tossing out hydoxide ion concentration arguments (sorry blogfast25) as per my previously cited "The Precipation of Aluminum Hydrous Oxides and its Solubility in Ammonia" (which I think is highly relevant here, may be some reader has full access and gives us some important historical research points) published in Analyst, Issue 767, 1940 by E. B. R. Prideaux and J. R. Henness. They 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)." 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.

I also wish to note another partial comment available on the 1st page "In the precipitation by ammonia, which depends on the almost complete hydrolysis of ammonium aluminate..." where the precipitation is most likely referring to Al(OH)3. Interestingly, this statement may imply the authors conditional acceptance of the existence of ammonia aluminate by the action of NH4OH on alumina.

However, as to the answer to my question, I am not clear as to your answer.

LanthanumK - 22-6-2011 at 12:22

Add one Cl to H9SCl4O3 and get the "compound" H9SCl5O3, which is a conglomerate of three components found in the reaction of sulfur dichloride with water: SCl2, 3 HCl, and 3 H2O. :P

Neil - 22-6-2011 at 12:26

As a simplistic experiment I heated some Al foil with an aspirated butane torch using as oxidizing of a flame as I could. The Al foil was then crumpled and pushed into a test tube filled with distilled water and swirled lightly until all visibly entrapped air was removed, I added a small amount of purple cabbage juice so that the burnt Al was sitting in a light purple liquid.

Given the sensitivity of cabbage juice and it's ease of acquisition this seemed like the simplest way to test for hydrolysis of AlN. No result, I heated the test tube after three hours to see if that would have an effect but still change in pH of gas production.



I added a tiny amount of NaCl to the test tube. I observed the very slow evolution of gas. The test tube was allowed to sit for two hours with no increases in the rate of gas production or changes in pH.

After two hours the test tube was set in a beaker of boiling water which substantially increased the rate the gas was produced. Gas was collected and ignited, gas did not produce "smoke" when exposed to HCl vapors.


So, the same results but with no Ammonia, and no change in pH.


Off topic but on subject-
This wouldn't sound familiar to anyone, would it?

"However, if you are a chemistry student or answering a question on an AP Chem exam, please ignore as I would seriously doubt the state of knowledge on real world chemistry among our educators."

AJKOER - 22-6-2011 at 12:53

Neil, even though I suspect you are not presenting a serious theory, believe it or not, a time study of the action of AlN in acidified water would be valuable, especially, if the Al completely dissolves in a few days.

However, this is off topic as the focus is not reactions involving AlN some much as involving the burnt Al with NH4OH + Salt.

blogfast25 - 22-6-2011 at 12:59

Quote: Originally posted by AJKOER  
Neil thanks for the picture.

Talk about trashing theories, I think we can start with tossing out hydoxide ion concentration arguments (sorry blogfast25) [...]


Toss away for all I care. The low alkalinity of NH3 solutions explains perfectly why alumina, hydrated or not, is essentially insoluble in those solutions. Any secondary effects will lead at best to very small solubilities.

Just because you make a 1940 paper the latest centre of your bizarrerie doesn't make it right or relevant. If hydrated alumina forms small amounts of a collodial Al(OH)3 in the presence of of NH3 this has nothing to do with 'ammonium aluminate'.

Neil, I doubt very much if AlN can be obtained in these conditions. If it could (and assuming it hydrolyses quickly - only an assumption right now) we'd all be using it as 'portable NH3'...

[Edited on 22-6-2011 by blogfast25]

[Edited on 22-6-2011 by blogfast25]

Neil - 22-6-2011 at 13:11

It is very on topic.

In removing variables to come to a correct conclusion one needs to... remove variables.


The reaction proceeds with ammonia, the reaction proceeds without ammonia.

It does not seem that ammonia is part of the reaction.

It doesn't get much simpler then that.

This is stupid. For a clearly failed 'one time chemistry major' you sure seem to have a very high opinion of your abilities. Reading a journal and then misapplying what you've read and then repeating it so often you think it is true IS NOT SCIENCE.

Did you have Salt in your experiments? If now yes, then you've lied before. If no, then you are an idiot for concluding that salty ammonia solutions are the same as salt-less ammonia solutions and that this in anyway applies to your conclusions.

If you can't see the relevance of a reaction taking place in spite of the loss of it's main proposed agent- then you are beyond daft.


Neil - 22-6-2011 at 13:17

Quote: Originally posted by blogfast25  

Neil, I doubt very much if AlN can be obtained in these conditions. If it could (and assuming it hydrolyses quickly - only an assumption right now) we'd all be using it as 'portable NH3'...




I 110% agree.:D

From what I've read it would be almost impossible for measurable amount to form, I was trying to devise an easily reproducible way to verify this to test the wild notion that it was relevant. Whatever the case, I give up.

AJKOER - 22-6-2011 at 13:42

Sorry to disagree, but there is still ammonia!

One of the most understated and ignore reaction in this whole thread is:

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

From the Salt Cake reaction slides, no-one is adding NH3, it is all from the AlN, and the reaction proceeds exothermically. All the energy needed to produce the AlN is released in its decomposition resulting in fires from the presence of H2 (no one mentioned the energy supplied by AlN, only the energy required to split Al2O3).

So sorry to break up the party, go back to the slides!

Neil - 22-6-2011 at 14:25

Quote: Originally posted by AJKOER  
Sorry to disagree, but there is still ammonia!

One of the most understated and ignore reaction in this whole thread is:

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

From the Salt Cake reaction slides, no-one is adding NH3, it is all from the AlN, and the reaction proceeds exothermically. All the energy needed to produce the AlN is released in its decomposition resulting in fires from the presence of H2 (no one mentioned the energy supplied by AlN, only the energy required to split Al2O3).

So sorry to break up the party, go back to the slides!





You're a dumb ass. No pH change = No dissolved ammonia

No reaction of the gas with HCl vapors = No ammonia gas

Likelihood of AlN formation? Almost zilich.

AJKOER - 22-6-2011 at 15:24

It is amazing how quickly you performed your undisputable experiment and got such results, it takes days for me to dissolve burnt Al (you minutes in weak acid, impressive!). By the way, in your cabbage juice, any fermentation (CO2)?

To quote from the burning of Al paper supplied in my 2nd post:

"After the second stage AlN content exceeded 50 mass% and residual Al content decreased to 10 mass%."

To an open mind, the Salt Cake example may or may not be extreme example. Why, because the AlN is tightly bound to both the Al2O3 and Al. Its removal upon reacting with water, may cause exposure to the Al and put fresh NH3 in contact with Al2O3. Remember, this is a slow reaction.

There may also be other chemical pathways, for example, assume the presence of CO2, then NaCl + NH3 + CO2 produces NaHCO3 and most importantly NH4Cl, which would produce havoc on the Al2O3.

I am telling these to you because if you insist on providing odd ball experiments, (rather than admitting that the plausible products to a question you were invited to answer, are all unacceptable to your position), please avoid experiments that can be explained by these plausible reactions.



[Edited on 22-6-2011 by AJKOER]

[Edited on 22-6-2011 by AJKOER]

[Edited on 23-6-2011 by AJKOER]

blogfast25 - 23-6-2011 at 05:33

Quote: Originally posted by AJKOER  
To quote from the burning of Al paper supplied in my 2nd post:

"After the second stage AlN content exceeded 50 mass% and residual Al content decreased to 10 mass%."

To an open mind, [...]


I hope you don’t seriously believe you’ve produced 50 % AlN with your Al ‘burning’. If so, prove it or at least provide evidence. In strong, hot acids a mix of Al2O3/AlN should probably digest quite quickly and release ammonia. And there must be other fairly simple ways to demonstrate the presence of significant (macroscopic quantities) of AlN.

It was only a matter of time before ‘open mindedness’ came marching with ill-deserved confidence towards this conversation. There are some things we hold proved beyond reasonable doubt: I don’t keep an open mind with regards the shape of the Earth for instance. You, on the other hand, have in this thread shown a tendency to propose reaction mechanisms that we KNOW cannot proceed. Science isn’t about reinventing the wheel at every turn in the road.

Perhaps universal indicator paper would have been a better choice than cabbage juice but for it to ferment you need sugar, yeast, the right temperature and quite a bit of time: it’s therefore unlikely Neil observed fermentation CO2.

It’s also quite reasonable to assume that small amounts of AlN, even at RT in contact with pure water, would affect the pH quite quickly: that’s kind of the nature of the water equilibrium (2 H2O < === > H3O+ + OH-); small amounts of even weak acid or base can affect it measurably (by simple means).

[Edited on 23-6-2011 by blogfast25]

AJKOER - 23-6-2011 at 06:57

Good point, the only counter argument is the the Al and Al2O3 are inherently amphoteric (acting here as a weak acid) and in the presence of a dilute weak base (NH4OH), the pH change may be moderated.

Also found a reference on the hydrolysis of AlN, 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."

Also found a full paper discussing the reactions in salt cake at: https://netfiles.uiuc.edu/tstark/website/Journal_Papers/JP84...

Per the paper "Aluminum Waste Indicators in an MSW Landfill", a possible important point is that the formation of the Al(OH)4 complex starts to occur when pH rises above 9 (note, the pH of household ammonia is generally given as over 11). Later in the article it is also noted that CaO reaction with water is also contributing to the rise in the pH.

I will let others read deeper and assess the significance of this work.

[Edited on 23-6-2011 by AJKOER]

[Edited on 23-6-2011 by AJKOER]

blogfast25 - 23-6-2011 at 13:24

When I think about it, the chances of creating any AlN by burning (oxidation with air oxygen) aluminium are pretty slim. Al2O3 has a heat of formation (HoF) of about - 1,676 kJ/mol (NIST value), NIST database lists AlN as having a HoF of a mere -318 kJ/mol (surprisingly small actually). Heating Al in a nitrogen/oxygen mixture would tremendously favour the formation of aluminium oxide over aluminium nitride.

In fact any formed AlN would likely burn to Al2O3: AlN + 3/2 O2 === > 1/2 Al2O3 + ½ N2, a reaction that’s highly exothermic (HoR = - 520 kJ/mol of AlN, if the HoF NIST value of AlN is to be believed).

AJKOER - 24-6-2011 at 14:21

I believe the actual amount of AlN formed is best left to direct measurements or prior studies. My logic is that there may be reactant sequences not fully understood in our procedures, and then there are the many kinetic and thermodynamic variables including varying temperature zones, heating times, Aluminum particle size (namely foil vs. powder) and air flow rate issues.

As a path and zone example, say my turbulent Al burning of methane in air first produces Al2O3, but then in my heating procedure there could be a methane reduction reaction occurring in a lower temperature zone (like inside the methane flame). So:

3 Al2O3 + 3 CH4 --> 3 CO + 6 H2O + 6 Al
(Correction below per Neil's reference at 1498 C:
Al2O3 + 3 CH4 --> 3 CO + 6 H2 + 2 Al )

Now, it may be that the Al2O3 is more subject to reduction by methane (or CO) than AlN.

Also, as you noted, in zones where the temperature exceeds 700 C, the AlN may be reduced by O2, but not the Al2O3:

4 AlN + 3 O2 --> 2 Al2O3 + 2 N2





[Edited on 24-6-2011 by AJKOER]

Neil - 24-6-2011 at 15:19

Quote: Originally posted by AJKOER  

As a path and zone example, say my turbulent Al burning of methane in air first produces Al2O3, but then in my heating procedure there could be a methane reduction reaction occurring in a lower temperature zone (like inside the methane flame). So:

3 Al2O3 + 3 CH4 --> 3 CO + 6 H2O + 6 Al

Now, it may be that the Al2O3 is more subject to reduction by methane (or CO) than AlN.

Also, as you noted, in zones where the temperature exceeds 700 C, the AlN may be reduced by O2, but not the Al2O3:

4 AlN + 3 O2 --> 2 Al2O3 + 2 N2


:o

http://opensourceecology.org/w/images/c/c9/Halman.pdf

blogfast25 - 25-6-2011 at 05:44

Quote: Originally posted by AJKOER  
My logic is that there may be reactant sequences not fully understood in our procedures, and then there are the many kinetic and thermodynamic variables including varying temperature zones, heating times, Aluminum particle size (namely foil vs. powder) and air flow rate issues.

[Edited on 24-6-2011 by AJKOER]


To determine whether or not a particular reaction, no matter how complicated its actual path may be, is thermodynamically feasible, we need only to consider the overall reaction’s ΔG (because G is a state function), see Hess’ Law.

Tinkering with the actual path (e.g. use of catalysts) can make a thermodynamically favourable reaction actually feasible in practical conditions, but it cannot make an ‘impossible’ reaction possible.

Neil, that’s great article but the reaction conditions are very far removed from what can be achieved with Al foil and a blowtorch, never mind also the advanced analytical technique of XRF…

AJKOER - 25-6-2011 at 08:16

I agree that your calculations are consistent with the 1st stage burning of Al. To quote the abstract from "Study of aluminum nitride formation by superfine aluminum powder combustion in air" by Alexander Gromov and Vladimir Vereshchagina, "The products of the first stage consisted of unreacted aluminum (70 mass%) and amorphous oxides with trace of AlN." However, in the the second stage (which involved the self sustaining burning of Al) "AlN content exceeded 50 mass% and residual Al content decreased to 10 mass%." Now, my question is to what extent (if any) does our lower temperature burning of Al ever achieve the benefit of the second stage, whatever reaction paths that may be involved? Your position is perhaps not at all, I am a little more positive.

Now my previously cited reference of the hydrolysis of AlN is very interesting in that the induction period is eliminated in the presence of a base. In other words, the immediate bubbling reaction observed in the presence of NH4OH could be completely explained if there is significant AlN.

Of course, if there is no AlN, it is also possible that a similar removal of the induction period occurs for Al2O3 (given another reference that an Al(OH)4 complex starts forming when pH exceeds 9 in the presence of NH3 and Ca(OH)2).

As an interesting side note on thermodynamic arguments, to quote from D.M. Dickson paper "Alumina from Coal Waste through the Formation of Aluminum" (see link below) "However, thermodynamic data alone are not sufficient in predicting the reaction rate or path." Dickson goes on to note in the case of a gas formation in one of several thermodynamically qualifying reactions, the favored reaction is the one with the highest sustainable partial pressure for the gas created.

LINK (see page 19)
http://content.lib.utah.edu/cdm4/document.php?CISOROOT=/undt...
'


[Edited on 25-6-2011 by AJKOER]

blogfast25 - 25-6-2011 at 12:57

I’m deeply sceptical about the validity of the Russian study but it’s hard to judge without the full article.

Regarding Dickson and gasses, that’s true, is well known and often practiced: the removal of a gaseous reaction product can pull a thermodynamically unfavourable reaction to the right but that doesn’t violate any known Law, quite the opposite. Nothing new. Also, how does it sustain your claims?

I maintain that in all likelihood yours and Neil’s attempts at synthesising AlN are far removed from any ‘real world’ AlN synthesis and that you have yet to provide any actual, reproducible evidence for any actual AlN formation in your experiments.

AJKOER - 25-6-2011 at 14:16

I was thinking of rolling Al foil to form Al rods. Then, it may be easier (or not) to produce an obvious self-sustaining Al combustion. As Al when ignited burns at over 2,000 C, this may reproduce the self-sustaining temperature witness in the Russian study, which, interestingly, indicated a favored reaction for the formation of AlN over Al2O3.

With respect to thermodynamics, whether or not the influence of gases are pertinent here, I am not sure. There is, however, at about 1500 C, a possible reduction reaction on Al2O3 with CH4 producing gases as was noted previously. If one believes in the final results of the Russian study at all, then it is reasonable to suspect some reduction paths at work on the Al2O3. My comment, however, was intended to be more of a general cautionary nature.

Neil - 25-6-2011 at 15:09

I've been looking for a home route to AlN in testable amounts but am having no luck. If you check the papers you've posted regarding its intentional creation, you will find that the papers are describing the reaction of Al and N under a nitrogen atmosphere. In the Al in air papers, the reaction is inside a sealed bomb. The reaction first uses up the available oxygen and then proceeds to combine with the available nitrogen. In many of the papers out of China, they are using a nitrogen flush or a sealed nitrogen chamber.

Eg,
http://ir.lib.ncku.edu.tw/bitstream/987654321/99518/1/301030...

In a number of papers examining Aluminum combustion they do not even acknowledge the reaction of Al + N2 when oxygen is present and at times use N2 instead of Ar as a inert gas.
http://www.dtic.mil/cgi-bin/GetTRDoc?Location=U2&doc=Get...
http://www.aero.polimi.it/~merotto/dottorato/Alcomb.pdf
ftp://ftp.rta.nato.int/PubFullText/RTO/EN/RTO-EN-023/EN-023-...

Combustion species of Al with many delta H values.
http://www.eng.buffalo.edu/~swihart/Reprints/Swihart_CombFla...


My computer went down last night and I am on another so I don't have my sources at hand...

However:
AlN oxidizes to Aluminum oxide with great ease, so much so that it seems it is nearly impossible to form it in meaningful amounts without the wash of hot pure nitrogen. It has been speculated that the aluminum enters a gaseous phase for Al+ N the reaction to proceed, the use of chlorides was mentioned in possibly catalyzing this.

In searching for aluminum combustion studies which were not done in sealed bombs or under a inert atmosphere, I am unable to find any papers which affirm the creation of AlN in normal air/aluminum combustion.

Interestingly this reminds me of speculation with regards to classical thermite mechanisms and if they make use of a gas phase during their reaction. Chlorides have the effect of increasing the burn rates of thermite mixtures through perhaps similar mechanisms

http://www.scribd.com/doc/47751250/Thermite-Reactions-their-... –PDF of this is in the reference section.


Based on my current research; I'm considering a reaction of aluminum nitrate nonahydrate with urea and glucose to form a Al2O3/C powder as per http://www.sciencedirect.com/science/article/pii/S0025540807...

There are a number of papers which examine this route but the PDF copies I had are all on the currently down computer.

In the papers they detail using a nitrogen furnace to roast the powders, I wish to try nitrate ignition using a mixture of Al2O3/C/KNO3/NH4Cl/Al

To be honest, based on what I've read so far the chances of me being able to isolate even 1g of AlN is close to 0.



Now;
Regarding reaction rates, I never said I fully dissolved/degraded the burnt foils. Currently I still have the original foil/ammonia/KCl reaction sitting in a corner. It has stopped showing any signs of reacting and has developed a thin white film over the bottom of the container. The remaining burnt foil is a black/brown color now.

Quote:
I am telling these to you because if you insist on providing odd ball experiments, (rather than admitting that the plausible products to a question you were invited to answer, are all unacceptable to your position), please avoid experiments that can be explained by these plausible reactions.


I'm not certain how my experiments are oddball. I used a absurd indicator, true, but that was in effort to make the experiment re-producible so you could, ya know, confirm or deny the results on your own without the need to acquire anything that was not strictly OTC.

Currently I've looked for a reaction of burnt foil with house hold ammonia, the reaction of burnt foil with ammonia contaminated with a chloride, the reaction of burnt aluminum with a chloride solution which contained no ammonia, the reaction of non burnt foil with ammonia and potassium chloride and attempted to find evidence of AlN formation by looking for pH changes when burnt Al foil was left in water.

So far I've observed that the reaction of the foil proceeds with or without burning, so long as a fresh surface is exposed.
I've observed that the reaction proceeds with or without ammonia, so long as a chloride is present. I've contained and tested the gas produces by these reactions and found it was not ammonia, nitrogen or oxygen but seemingly hydrogen.
I've observed no gas formation or pH change which would be indicative of the hydrolysis of any formed AlN on burnt foil

Please, if these are oddball and irrelevant, explain how you would test your proposed theories and then please pretty please; run the experiments. Visa versa, currently you're proposing odd ball explanations with implausible reactions to fit your continually evolving position while offering no experimental data to validate your position.

In attempting to re-stock my ammonia supply I discovered the bottles of colorless/unscented Ammonia are no longer available. I'm having some luck with pulling the nasty yellow lemon scented Satan spit that is available, through charcoal. There is some off gassing I'm not sure how much... I need a burette...

Was your ammonia colorless and unscented?

With regards to burning Al foil, not going to happen. Even Al turnings, dampened with water and piled high only crumble under a air/MAPP flame

@ Gert Re Alumina reduction; Sorry, sarcasm is swallowed by text.

AJKOER - 25-6-2011 at 16:20

Good question on the NH4OH. Brought both a yellow and a clear in Pathmark. I performed some runs using the yellow NH4OH, no difference other than the world is viewed through a yellow lens.

I am planning on running some purified NH4OH experiments (to remove the influence of the detergents).

The AlN effect in our samples appears to be becoming more of a trace issue.

FYI, the reason I have focused on burned Al is to better understand what is attacking the protective Al2O3 (minus some of the coating defenses), I have also observed similar dissolving of Al (especially after pre-soaking in vinegar to weaken its defenses) in NH4OH.

[Edited on 26-6-2011 by AJKOER]

blogfast25 - 26-6-2011 at 05:09

Another thing that would favour oxide formation over nitride is when you carry it out under pressure: Al + 3/2 O2 === > ½ Al2O3
Al + ½ N2 === > AlN

Neglecting for a minute the molar volumes of oxide and nitride, the oxide forming reaction leads to a significantly higher reduction in volume from left to right. That’s a classical ‘forcing condition’ that’s often imposed, see e.g. N2 + 3 H2 < === > 2 NH3 at very high pressure (Born Haber process).

For reaction of Al with air, the situation is probably similar to Ca: Ca is a scavenger of nitrogen air but only if enough of it is available (and air supply is limited). The first part then gobbles up the oxygen, the remainder the nitrogen.

Any backyard science research into AlN should probably concentrate on trying to create such conditions. Some fine aluminium, a quartz tube, a couple of gas syringes, a Meker burner and some decent planning/calculations could do wonders there...



[Edited on 26-6-2011 by blogfast25]

AJKOER - 26-6-2011 at 05:50

Significance MSDS Results 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:...

As Aluminum is coated with Al2O3, it is meaningful to search Aluminum MSDSs as well. Some incompatibility statements with Al are consistent with experiments performed on this thread. In particular, the incompatibility of Al with NH4NO3, acid chlorides and "metal salts". Note the last reference is in agreement with the mineral content of Salt Cake and also could be referring to the known electrochemical properties of Al metal in aqueous salt solutions acting as the electrolyte with metals like Silver, for example. Also, there is a stated incompatibility of Al with water, and the reaction of water vapor and molten Al is reportedly explosive. As such, this water/Al reaction apparently cannot be attributed to dissolved salts.

Al MSDS Links:
http://www.sciencelab.com/xMSDS-Aluminum-9922844

http://www.floodbreak.com/default/Maintenance%20Ops/Aluminum...

[Edited on 26-6-2011 by AJKOER]

blogfast25 - 26-6-2011 at 09:43

Dear Flying Spaghetti Monster, AJKOER, resorting the MSDS sheets now, eh? Ridiculous…

The idea that ammonia is a catalyst for aluminate formation is ludicrous for one simple reason. The equilibrium:

NH3 + H2O < === > NH4+ + OH-

… in the presence of a strong alkali is completely pushed to the left, even by small amounts of strong alkali. This can be calculated very, very easily and quite accurately from the equilibrium constant K<sub>b</sub>, or, in a mildly more complex algebraic development from K<sub>b</sub>, K<sub>w</sub> (water dissociation constant), mass balances and solution neutrality requirement. It’s classic 3rd year university stuff.

Most ammonium salts (from strong acids) are aluminium incompatible because these salts react slightly acidic in watery solution:

NH4+ + H2O < === > NH3 + H3O+

As a weak acid, NH4+ is well understood. NH4NO3 solutions (but also halides, sulphates and other ammonium salts from strong acids) are slightly acidic, not recommended for use with a reactive metal like Al (passivation layer OR NOT). Simples.

AJKOER - 26-6-2011 at 09:45

Neil:

With respect to a possibly easier way to make AlN, see the extract below from the US Patent 5226952 which apparently employs a replacement reaction with calcium nitride and Al (and also a CaAl2 impurity) that forms AlN and pure Ca, the latter being the apparent goal of the process. Note, the formation range of 700 to 1200 C.

"United States Patent 5226952

A process for refining calcium containing aluminum as an impurity in the form CaAl2. The calcium is nitrided to form Ca3N2, and that Ca3N2 is made to react with CaAl2 initially present and with the particulate aluminum in such a manner as to form calcium, which is isolatable in the form of high-purity calcium containing less than 0.1% by weight of aluminum, and aluminum nitride.

7. The process of any one of claims 1-3, in which the calcium is nitrided with the aid of nitrogen at a temperature of between 200° and 350° C.

9. The process of one of claims 1-3, in which the calcium nitride formed is made to react with the particulate aluminum introduced and with CaAl2 initially present in the calcium, at a temperature of between 700° and 1200° C., and the calcium formed is separated by distillation.

11. The process of one of claims 1-3, in which a calcium to be purified, containing n mols of calcium and m moles of CaAl2, is nitrided in such a manner as to form n/3 mols of Ca3 N2, which is made to react with m mols of CaAl2 and (n-3m) 2/3mols of aluminum, so that aside from losses, n+m mols of purified calcium and 2n/3 mols of AlN are obtained. "

You may be able to purchase Ca3N2, and then react it with particulate Al between 700 to 1200 C, if able and so inclined. Note, there is a parallel employment of Calcium carbide to form carbides of other metals.

[Edited on 26-6-2011 by AJKOER]

AJKOER - 26-6-2011 at 10:12

blogfast25

Please read my post again.

I am not dictating any aluminate formation (perhaps you missed the word "not"), but I am suggesting an aqueous alkaline catalyst to the reaction of Al2O3 and water (with parallels to AlN hydolysis). The fact that NH4OH is predominately NH3 and most importantly H2O makes it a good "aqueous alkalie solution".

I also noted that this was not an original concept on my part.

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 your concerns on the quality of some MSDS statements.

Link:

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

[Edited on 26-6-2011 by AJKOER]

blogfast25 - 26-6-2011 at 13:16

My concern isn't with the 'quality of some MSDS statements' (although many are pretty crap), but with the fact that MSDS sheets aren't scientific literature, rarely specify any conditions etc. They can be OK to glean the odd specific property of a compound or product from but not much besides that. It's not really their function to provide precise chemical information.

How do solutions of unprotonated NH3 make "good aqueous alkaline solution"?

AJKOER - 26-6-2011 at 14:23

How do solutions of unprotonated NH3 make "good aqueous alkaline solution"?

I would guess pretty much the same way that it acts on AlN to initiate the hydrolysis action.

As far as I know, more details on the precise catalytic mechanism may, or may not, actually be known.

AJKOER - 26-6-2011 at 17:15

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:...

blogfast25 - 27-6-2011 at 05:10

If it’s unprotonated, it’s not an alkali: for something to qualify as alkali OH- ions must be present.

The paper is very interesting but glaringly makes my case for me: the amounts of dissolved/peptised/supersaturated hydrated alumina are very small and tail off with increased ammonia concentration, despite the fact that the OH- increases slowly with increasing NH3 concentration.

It confirms that using NH3 solution to precipitate hydrated alumina from Al<sup>3+</sup> bearing solutions is good practice and easier than using strong alkalis, where overshooting the required OH- concentration is often a problem.

AJKOER - 27-6-2011 at 19:09

Here is an explanation why NH4OH is the proper formula for ammonia in water:

"nitrogen likes to form 3 bonds to complete it valence shell, but it also has that lone pair of unbonded electrons hovering in that p orbital. In a water environment, nitrogen's electron orbital will actually dissociate water into H+ and OH-. The H+ is of course attracted to nirtogen's negatively charged electron cloud and it forms the ammonium ion, NH4+, which ionically bonds to the remaining OH- hydroxide ion from the water"

source: Kayak on Ask.com

barley81 - 27-6-2011 at 19:26

Yes, but in ammonia solutions, only a small amount of NH4+ is formed. In household ammonia, the pH is around 11-12 (~0.01M OH-). The great majority of ammonia exists as NH3.

blogfast25 - 28-6-2011 at 04:30

Quote: Originally posted by AJKOER  
Here is an explanation why NH4OH is the proper formula for ammonia in water:

"nitrogen likes to form 3 bonds to complete it valence shell, but it also has that lone pair of unbonded electrons hovering in that p orbital. In a water environment, nitrogen's electron orbital will actually dissociate water into H+ and OH-. The H+ is of course attracted to nirtogen's negatively charged electron cloud and it forms the ammonium ion, NH4+, which ionically bonds to the remaining OH- hydroxide ion from the water"

source: Kayak on Ask.com


Yeah, never forget to ask ‘kayak’ from ask.com, he’ll know!

It just shows how little you know about chemistry and how you’re scraping the barrel.

‘NH4OH’ is a very old way of notation of NH3 solution, dating back to before Bronsted acid-base theory.

NH3 does have an extra electron pair which it uses to bond protons (from water) with (that makes it also an electron donor). The reaction is thus as indicated ad nauseam above:

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

… which proceeds only feebly (the product of [NH4(+)] and [OH-], divided by [NH3] is only about 1.78 x 10<sup>-5</sup>;)

What little NH4(+) forms is completely dissociated from OH-. There never was any ‘NH4OH’ to begin with but once upon a time its believed existence was a reasonable explanation for ammonia’s weak alkalinity.

Sigh… :(

AJKOER - 28-6-2011 at 08:36

"Ammonia is not a base when using the more widely known Arrhenius definition (which states that a base releases OH- and an acid releases H+. Instead, it is considered a base when using the Lewis definition, which states that a base is an electron pair donor, and an acid is an electron pair acceptor. The structure of ammonia, NH3, has a pair on the nitrogen atom, and it is this lone pair that acts as a electron pair donor.

Also, when mixed with water, ammonia forms ammonium hydroxide (NH4OH), which is a base according to the Arrhenius definition."

Link:
http://wiki.answers.com/Q/Why_does_NH3_give_OH-_ions_on_diss...


Also here is a reference from "Chemistry Explained", a teaching website:

Ammonia (NH 3 ) is a weak base, and although it does not have OH − ions in its formula, it produces the ion on reaction with water.

NH 3 ( aq ) + H 2 O ( l ) ⇆ NH 4 + ( aq ) + OH − ( aq ) (8)

Link:
http://www.chemistryexplained.com/A-Ar/Acid-Base-Chemistry.h...

Apparently, this view is generally held, for example:

"Similarly, ammonia dissolves in water and partly dissociates in the presence of water into ammonium ions. As we all know that water dissociates into H(+) and OH(-). The NH3 reacts with H(+)to form NH4(+). So we can write a chemical reaction:

NH3 + H(+) + OH(-) <-----> NH4(+) + OH(-)

The above reaction is an equilibrium reaction. So the amount reacts depends on the equilibrium constant which is dependent on temperature."

Link:
http://www.cheresources.com/invision/topic/9750-ammonia-wate...

While all interesting, the argument on NH4OH is not relevant since NH4OH is generally considered a weak base and I believe this is what was precisely being reference on the MSDS and you, I suspect, do not.

blogfast25 - 28-6-2011 at 09:08

Widely held views aren’t necessarily true. Truth isn’t arrived at by majority rule.

Teaching sites simplify things accordingly. My daughter passed (I hope) her GCSE chemistry exam today: on that level they’re still writing reaction equations in the simplest of forms. Actual chemists really only do so when writing overall reactions or for stoichiometry determination. And the definition of ‘catalyst’ she was taught is comical to someone like me…

Water based solutions contain almost no free protons (H<sup>+</sup>;): these are heavily solvated to H<sub>3</sub>O<sup>+</sup> (hydronium or oxonium ions).

NH3 is both a Lewis AND a Bronsted base. Arrhenius acid-base theory is now largely obsolete.

The question ‘why does NH3 give OH- when dissociated in water’ is badly formulated. NH3 isn’t ‘dissociated’: it’s (weakly) protonated in water. The every small proportion of water molecules thus deprotonated become OH- ions.

Do you think you could go and read quietly now?

AJKOER - 29-6-2011 at 06:50

blogfast25:

Another point is I agree with your opinion that any alleged ammonium aluminate is best described (to quote you) as "dissolved/peptised/supersaturated hydrated alumina".

However, a subtle legalistic and possibly important point is that there is an alluded to process by which ammonium aluminate has been prepared, namely "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." Reference: "Journal of the American Chemical Society", Volume 38, page 1287 (link previously supplied).

The potential legal significance of this is that to quote from 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, I have seen ammonium aluminate cited in a patent, possibly since one cannot claim its "indefiniteness".

Link: http://www.uspto.gov/web/offices/pac/mpep/documents/2100_217...

blogfast25 - 2-7-2011 at 06:37

Regarding AlN, Brauer’s ‘Preparative Inorganic Chemistry’ gives three methods for its preparation (p.827, available in SM’s library):

Aluminum
A1N
Nitride
= A1N
41.0
A nickel boat is filled with very pure aluminum powder which
has been degreased and dried either by extraction with ether or
by heating to 150°C in a stream of nitrogen. The boat is placed
in a quartz or porcelain tube and heated in an electric furnace
while purified nitrogen is passed over it. Even though the nitride
starts to form on the surface below 650°C, the reaction proper
begins only at 820°C, when the entire mass begins to glow. At
this point the flow of nitrogen should be increased to prevent the
Ns pressure from decreasing owing to the rapid reaction. When
the reaction is essentially complete, the mass is allowed to cool
in a stream of nitrogen. Since the product still contains some
unreacted metal, it is pulverized and reheated under nitrogen for
1-2 hours at 1100-1200°C. The product obtained is nearly white
and has a nitrogen content not far below theoretical.
II. Al + NH3 = A1N + 3/2H2
27.0 17.0 41.0
To obtain silicon-free A1N, aluminum powder pretreated as
above is placed in a trough of molybdenum sheet inside a nickel
reaction tube, and NH3 is led through while the tube is heated to
1300°C in an electric furnace.
III. AlCls • NH3 = A1N + 3 HC1
150.4 41.0 109.4
The reaction is performed in the apparatus shown in Fig. 246,
which consists essentially of a thick-wall Pyrex tube with an
enlargement in the middle and four necks at the top. A thin
glass tube (nitrogen inlet) passes through the middle neck and
reaches nearly to the bottom. The two side necks contain silver
wire leads to a tungsten heating coil w suspended in the reaction
tube at the level of the bulb. The fourth neck is an outlet for the
gas. The reaction tube is thoroughly dried and A1C13- NH3, prepared

Neil - 7-7-2011 at 16:30

Ah, that would make my plans a fools errand. Thank you for that :)

blogfast25 - 8-7-2011 at 06:13

Welcome!