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AJKOER
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Reaction between NH4OH and Al2O3
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]
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AJKOER
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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.
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bbartlog
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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.
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blogfast25
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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!
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LanthanumK
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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.
hibernating...
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blogfast25
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| 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.
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LanthanumK
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You did state it scientifically, but with overemphasis.
hibernating...
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blogfast25
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Sometimes overemphasis is good.
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AJKOER
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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]
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m1tanker78
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Ajoker: What method are you using to burn the Al foil? Seems easy enough to try and reproduce...
Tank
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sternman318
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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?
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AJKOER
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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]
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LanthanumK
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NaOH is not all that hard to make.
hibernating...
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redox
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How would you make it, without contamination by NaClO, NaClO4, and NaCl?
My quite small but growing Youtube Channel: http://www.youtube.com/user/RealChemLabs
Newest video: Synthesis of Chloroform
The difference between chemists and chemical engineers: Chemists use test tubes, chemical engineers use buckets. |
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blogfast25
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| 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.
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!
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m1tanker78
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| 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
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blogfast25
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| 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]
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m1tanker78
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| 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
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blogfast25
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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)
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LanthanumK
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If anything, it increases its solubility as the equilibrium NH4+ + OH- <=> NH3 + H2O is moved to the left.
hibernating...
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Neil
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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?
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AJKOER
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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]
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blogfast25
International Hazard
Posts: 10562
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1916, fantastic vintage! Next we’ll be reviving Phlogiston Theory to try and explain the basically non-existent ammonium aluminate.
| 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.
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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]
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Neil
National Hazard
Posts: 556
Registered: 19-3-2008
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| 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.
| 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.
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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...
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AJKOER
Radically Dubious
Posts: 3026
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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]
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