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blogfast25
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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?
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AJKOER
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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.
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LanthanumK
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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. 
hibernating...
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Neil
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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."
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AJKOER
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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.
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blogfast25
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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) [...]
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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]
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Neil
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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.
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Neil
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| 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'...
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I 110% agree.
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.
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AJKOER
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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!
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Neil
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| 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!
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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.
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AJKOER
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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]
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blogfast25
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| 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]
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AJKOER
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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]
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blogfast25
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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).
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AJKOER
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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]
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Neil
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| 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
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http://opensourceecology.org/w/images/c/c9/Halman.pdf
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blogfast25
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| 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…
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AJKOER
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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]
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blogfast25
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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.
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AJKOER
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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.
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Neil
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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.
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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.
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AJKOER
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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]
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blogfast25
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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]
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AJKOER
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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]
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blogfast25
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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.
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