AJKOER
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Solution of Dissolving Alumina in salt and ammonia proposed
I reported awhile back on my personal experience with the action of aqueous ammonia and salt on Al2O3/Al (there appears to be some bubbling). More
interesting I found an 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 (link: http://www.es.anl.gov/Energy_systems/docs/process_tech/indus... ):
First, the author's definition of Salt cake, to quote:
"Salt slag (Salt cake)
•By-product of aluminum recycling–Dross, post-consumer scrap, UBC, prime scrap •Salt cake composition: source material and recycling practice
dependent–Aluminum metal (finely distributed) (4-8%)–Salt (flux --NaCl, KCl, Fluoride) (25-45%)–NMP (non-metallic product) (50-70%)
•Oxides (mainly alumina and spinel–MgAl2O4) •Unrecovered aluminum
•Fluorides
•AlN(Aluminum nitride)
•Other oxides (Fe, Si, Zn..), carbides…"
Apparently, the author contends that via the reaction of Aluminum nitride with water:
AlN + 3 H2O ----) Al(OH)3 + NH3
the ammonia produced 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"
I now pivot from chemistry to electro-chemistry with some observations from the so-called Aluminum air battery as a possible explanation as to how
alumina becomes expose to a high pH so as to dissolve. To quote from Wikipedia (http://en.m.wikipedia.org/wiki/Aluminum_air_battery):
"Technical problems remain to be solved to make Al–air batteries suitable for electric vehicles. Anodes made of pure aluminium are corroded by the
electrolyte, so the aluminium is usually alloyed with tin or other elements. The hydrated alumina that is created by the cell reaction forms a
gel-like substance at the anode and reduces the electricity output. This is an issue being addressed in the development work on Al–air cells. For
example, additives that form the alumina as a powder rather than a gel have been developed.
Modern air cathodes consist of a reactive layer of carbon with a nickel-grid current collector, a catalyst (e.g., cobalt), and a porous hydrophobic
PTFE film that prevents electrolyte leakage. The oxygen in the air passes through the PTFE then reacts with the water to create hydroxide ions. These
cathodes work well but they can be expensive.
Traditional Al–air batteries had a limited shelf life[8] because the aluminium reacted with the electrolyte and produced hydrogen when the battery
was not in use–although this is no longer the case with modern designs. The problem can be avoided by storing the electrolyte in a tank outside the
battery and transferring it to the battery when it is required for use.
These batteries can be used, for example, as reserve batteries in telephone exchanges and as backup power sources. Al–air batteries could be used to
power laptop computers and cell phones and are being developed for such use."
And also:
"Aluminium based batteries
Different types of aluminium batteries have been investigated:
Aluminium-chlorine battery was patented by United States Air Force in the 1970s and designed mostly for military applications. They use aluminium
anodes and chlorine on graphite substrate cathodes. Required elevated temperatures to be operational.
Aluminium-sulfur battery worked on by American researchers with great claims, although it seems that they are still far from mass production. It is
unknown as to whether they are rechargeable.
Al–Fe–O, Al–Cu–O and Al–Fe–OH batteries were proposed by some researchers for military hybrid vehicles. Corresponding practical energy
densities claimed are 455, 440, and 380 Wh/kg[9]
Al-MnO manganese dioxide battery using acidic electrolyte. Produces a high voltage of 1.9 volts. Another variation uses a base (potassium hydroxide)
as the anolyte and sulfuric acid as the catholyte. The two parts being separated by a slightly permeable film to avoid mixing of the electrolyte in
both half cells. This configuration gives a high voltage of 2.6–2.85 volts.
Al-Glass system. As reported in an Italian patent by Baiocchi [[10]], in the interface between common silica glass and aluminium foil (no other
components are required) at a temperature near the melting point of the metal, an electric voltage is generated with an electric current passing
through when the system is closed onto a resistive load. The phenomenon was first observed by Baiocchi and after Dell'Era et Al.[11] began the study
and the characterization of this electrochemical system."
The actual half reactions occurring are given by (see https://www.google.com/url?sa=t&source=web&rct=j&... ):
"Anode: Al(s) + 3OH−(aq) → Al(OH)3(s) + 3e−
Cathode: O2(g) + 2H2O(l) + 4e− → 4OH−(aq)
Overall: 4Al(s) + 3O2(g) + 6H2O(l) → 4Al(OH)3(s) "
Now clearly, the reaction at the cathode can raise pH so as to enable the dissolution of the alumina (as an aluminate complex?) and permit the anode
half reaction as well.
--------------------------------
I created an aqueous version of the battery using Al foil, household dilute ammonia, dilute 3% H2O2 in place of air to speed things up and some sea
salt, a good elecrolyte. What is immediately visible are tiny bubbles (Hydrogen, I suspect). What was unexpected is how in 5 minutes the shiny
Aluminum foil becomes black (Silica or Carbon from the Aluminum alloy used to create the foil?)). White floating particles of Al2O3 are also apparent.
After a day, the reaction cannot be restarted by adding more H2O2 (Al anode is apparently coated as noted in commentary above).
What I am still unclear on is how the reaction starts with just weak ammonia and salt (which also serves as the electrolyte for the galvanic cell)?
Does adding salts increase the 'activity effect' (see general discussion of the topic at http://chemwiki.ucdavis.edu/Analytical_Chemistry/Analytical_... ) of the weak base so as to enable the jump starting of the reaction? Are
complexes being formed as side reactions to commence the half reaction at the anode?
Have I presented a better explanation as to what is being observed then my undoubtedly illustrious author at the Global Symposium?
[Edited on 23-6-2014 by AJKOER]
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DraconicAcid
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Quote: Originally posted by AJKOER  | | I created an aqueous version of the battery using Al foil, household dilute ammonia, dilute 3% H2O2 in place of air to speed things up and some sea
salt, a good elecrolyte. What is immediately visible are tiny bubbles (Hydrogen, I suspect). What was unexpected is how in 5 minutes
the shiny Aluminum foil becomes black (Silica or Carbon from the Aluminum alloy used to create the foil?)). White floating particles of Al2O3 are also
apparent. |
(Bolding mine). The bubbles are undoubtedly oxygen from the decomposition of hydrogen peroxide. You are not going to generate hydrogen gas in the
presence of a peroxide.
Please remember: "Filtrate" is not a verb.
Write up your lab reports the way your instructor wants them, not the way your ex-instructor wants them.
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AJKOER
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OK, the significant observed chemical side reactions that occurs with Salt cake is indeed the release of H2, per the reaction cited by the author. To
quote:
"dissolves the alumina film on unrecovered Aluminum particles surface, thereby exposing the Al surface to the reaction Al + H2O --> Al(OH)3 + H2"
There isn't must doubt on this as there are apparently fires from the burning Hydrogen at the Salt cake sites. The bubbles, I personally observed, are
interestingly micro and insignificant in total amount of H2 formed (well, at least, in my small scale experiment and not the bigger world).
However, I guess, with impurities present in the Al alloy, especially Fe forming Fe2O3, some decomposition of H2O2 to O2 is possible.
I thank you for your comment as it has just occurred to me that this Al/Hydrogen formed gas may be capable of reducing an oxide. See, for example, an
interesting 2012 published article on such active hydrogen, link http://www.researchgate.net/publication/221934434_Chemical_r... , titled "Chemical reduction of an aqueous suspension of graphene oxide by nascent
hydrogen" by Viet Hung Pham, Hai Dinh Pham, ... in Journal of Materials Chemistry (Impact Factor: 5.97). 05/2012; DOI:10.1039/C2JM30562C .
I doubt if this contributing to the removal of the Al2O3 coating, but a thought, similar to a recent cited article in the Journal of Materials
Chemistry per the source link.
I am increasingly convinced that this thread doesn't really belong in 'Beginnings'. For example, I have as of yet avoided discussing the literature on
the formation of complexes. More specifically, in the paper "The Precipation of Aluminum Hydrous Oxides and its Solubility in Ammonia" 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.
Yes, a lot of interesting material, but not really elementary.
[Edited on 23-6-2014 by AJKOER]
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MrHomeScientist
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Why don't you test the gas you produced, rather than just presuming? Gas tests for hydrogen and oxygen are quite easy to perform. It's almost
certainly oxygen, as chloride is a known catalyst for peroxide decomposition. Also, have you measured if your "battery" actually generates any
current? Stick a voltmeter across the terminals and see what you get.
Testing your setup will be very valuable to getting your questions answered, and is much more scientific than armchair theorizing - especially when
you have the experiment right in front of you!
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AJKOER
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Today I report the results of substituting for weak aqueous amonia a concentrated Sodium carbonate solution. Normally, the reaction of Al and pure
aqueous Na2CO3 proceeds only on boiling and quickly ceases on the removal of heat. However, a hot (not boiling) solution of Na2CO3, H2O2 and sea salt
is capable also of producing a self-sustaining reaction as does ammonia without heating.
--------------------------------------
A possible answer to the mechanics of my jump start question on the apparent electrochemical reaction. I can across an excellent white paper (link https://www.google.com/url?sa=t&source=web&rct=j&... ) detailing reaction of Aluminum and water with promoters to address the
protective Al2O3 layer. To quote from page 7:
"It has been shown that mixtures of aluminum and aluminum oxide (Al2O3) powders are reactive with water in the pH range of 4-9 (11-13) and at
temperatures of 10-90 oC. These Al-Al2O3 powder mixtures must be heavily ball-milled together in order to produce hydrogen reactions. Hydrogen can
be evolved at room temperature using essentially neutral water, although the hydrogen evolution rate increases with increasing temperature. "
Also, to quote page 8: "The aluminum oxide may be in the form of bayerite (Al(OH)3), boehmite (AlO(OH)), gamma alumina (γ-Al2O3), or alpha alumina
(α-Al2O3). Alpha alumina powder was reported to give the maximum hydrogen evolution. It has been speculated that the milling of aluminum and aluminum
oxide powders together helps to mechanically disrupt the adherent and coherent oxide layers present on the aluminum powder, and that this is the
reason for the enhanced hydrogen generation in pH neutral water (11-13).
However, recent research has suggested that the enhancing effect of aluminum oxide on the reactivity of aluminum with water may also be
mechanochemical in nature (14). Aluminum powders that were reacted with fine boehmite powders at elevated temperatures produced a layer of
fine-grained, mechanically weak gamma alumina on the surfaces of the aluminum powders. "
Apparently weakened gamma alumina reacts as follows:
Induction Stage: Al2O3 + H2O ---) 2 AlOOH
Followed by:
6 AlOOH + 2 Al ----) 4 Al2O3 + 3 H2
leading to the rupturing of Aluminum oxide layer.
So, the burning of processed Al foil could produce some weaken gamma alumina as could occur also in the salt cake example. Given the moderate pH range
upon which the composition of Al and weakened gamma Al2O3 is subject to rupture per the above reactions, we have a tenative explanation of the
initialization reaction in the electrochemical setting.
[Edited on 23-6-2014 by AJKOER]
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aga
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Seeing as this Is in beginnings, can you please post a quick experiment for the Beginners here, namely me ?
If you can say what the concentrations and quantities of the reactants are, and what the reaction procedure is, i'll do the experiment.
Have several reagents to test the outcome handy.
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AJKOER
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MrHomeScientist:
I did attempt to test for hydrogen gas, but not surprisingly based on observations of how little gas was being produced, it was not conclusive.
The fact that so little gas is evolving is in support of the electrochemical depiction as per the net reaction:
4Al(s) + 3O2(g) + 6H2O(l) → 4Al(OH)3(s)
no H2 is formed, except apparently per a competing side reaction of Al with water.
Other observations include that no reaction proceeds without salt (an electrolyte is required), the reaction restarts if depleted H2O2 is added, and
most telling, the Al does not completely dissolve as when all of the Aluminum is coated, the reaction just stops, but adding fresh Al, it restarts.
I strongly recommend others test out the galvanic cell and even experiment on replacing the ammona with other weak bases (as I did using Na2CO3).
----------------------------------
With respect to quantities to employ, the beauty of an electrical reaction is that it is usually of little consequence. Too little Al, and you can see
that you need more as it is consumed. When the H2O2 is consumed, but there is still air exposure, the reaction will really slow down. The ammonia is a
catalyst to provide a pH within an acceptable range. If the pH is off, the reaction will not start or stops. Pretty forgiving in general.
[Edited on 24-6-2014 by AJKOER]
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aga
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OK. Will do, if you would please specify how you want it doing.
e.g.
1g NaCl
2g Al foil cut into bits
beaker
2 nails as electrodes
100ml water
2ml 5.6[M] ammonia
or similar ...
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AJKOER
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Aga:
No nails (we don't want an iron air battery). You already have a cathodic zone, your Al flakes, or a single stick of Al, which, however, I would not
recommend as this would have very limited surface area, and the Al air battery has cathode coating issues. A single sheet of Al foil perhaps.
The anode is also a zone, the aqueous H2O2 and/or air contact (I recall a video on the Aluminum air battery that uses porous carbon to capture O2 if
you really want to hook up wires, and NaOH or KOH as the electrolyte in that battery).
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aga
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so ...
post the experiment you want done, and how to do it then ?
It is not difficut to do.
Just type 'get this, this and this, put it together like this, and then do this'
Substitute 'this' for the substances, concentrations, volumes/masses and the prcess you want following.
Then it can actualy get done.
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AJKOER
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My most recent experiment for those interested:
Following the chemistry I posted above, I first heated to glowing a piece of Al foil 12 inches x 6 inches. I left a few small areas unburnt ( shiny)
for visible observations. I added the Al/Al2O3 to a glass vessel, followed by 20 cc of sea salt (it was loaded with moisture /caking) and 60 cc of
dilute clear household ammonia. I swirl the contents for 30 seconds before I added 60 cc of 3% H2O2. I warmed the mixture by placing the lightly
sealed vessel in a microwave for 20 seconds (or, one could place the vessel in boiling water for a few minutes ). Done.
Within 15 minutes you should see obvious movement of Al/Al2O3 pieces as a result of attached bubbles. Some of the remaining shiny Al foil pieces are
becoming dark as well.
[Edited on 23-6-2014 by AJKOER]
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aga
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Yay !
Finally something i can duplicate !
Cheers.
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blogfast25
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Is this worth duplicating? That shouldn't be too hard: basically nothing happens. Talk about a damp squib... 
[Edited on 24-6-2014 by blogfast25]
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aga
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OK.Done that.
No idea what your 'black' is, unless it's Soot. Did you heat the foil with a candle ?
Five 1" square pieces of aluminium foil were heated to glowing red with a Butane pencil torch.
No blackening was observed.
These were placed into 5 beakers :-
1. Control, to which was added 20ml DIW
2. NaCl test, to which was added just 20ml of DIW.
3. KCl test, to which was added just 20ml of DIW.
4. OTC NH3 3w% test, to which was added 20ml of household ammonia + 10ml OTC 3w% peroxide.
5. Conc NH3 test, to which was added 20ml of 7.3[M] ammonia + 10ml OTC 3w% peroxide
After 15 minutes, samples 1, 2 and 3 showed no sign of change or activity.
Samples 4 and 5 were showing tiny bubbles evolving very slowly, when viewed under strong magnification.
At 25 minutes, 1g of NaCl(s) was added to samples 2, 4 and 5.
1g KCl was added to sample 3.
No immediate change was observed in any sample.
At 30 minutes samples 1, 2 and 3 showed no change.
Samples 4 and 5 were showing a colour change in the aluminium, with it going darker in patches.
Sample 4 showed more continuous bubble generation, albeit slowly, possibly more obvious due to the soap in OTC ammonia.
At 40 minutes samples 2 and 3 showed very slight bubble formation when magnified, with no other changes.
Samples 4 and 5 were breaking up into grey pieces, still with slow bubble generation. Approx 10% of the metal foil was still recognisable as such,
more so in sample 4.
Sample 1 showed no change or activity at any time.
In samples 4 and 5 the liquid had become cloudy.
Each sample was subjected to a test for Aluminium ions at 50 minutes.
1. Negative
2. Positive (highest reading)
3. Negative
4. Marginal. if positive, extreemly low reading
5. Positive
The bubbles formed too slowly to be captured for testing.
I can't remember the Point of this experiment, but enjoyed doing it.
Draw your own conclusions.
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AJKOER
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OK, some clarification on my perception on what is occurring may be in order.
First, preparing for the attack of the Al2O3 layer. To accomplish this ones burns household Al foil (I just repeated this, and to improve the process,
I move the lid on the stove burner slightly off which creates a louder so called perturbed methane flame which I read is over a 100 C hotter than an
unperturbed flame). My large sheet of Al foil was easily made red and I reheated as well. I then added to this Al/Al2O3 60 cc of clear dilute
household ammonia and 30 cc of a salt mix (25cc sea salt and 5 cc plain iodized NaCl). The reason for the excess salt is to increase the so called
'activity coefficient' (see my prior educational reference) and the effective strength of the aqueous ammonia. I then heated the mix for 20 seconds
with a loose cover in the microwave. To my shock, looking closely at the mix showed significant acivity (production of steams of a white compound and
bubbling). Note, I did NOT as of yet add any H2O2!
Apparently, I managed to create some weakened gamma Al2O3 bounded to Al, which per my cited previous alluded to white paper, is attacked up to pH 9
via the steps:
Induction Stage: Al2O3 + H2O ---) 2 AlOOH
Followed by:
6 AlOOH + 2 Al ----) 4 Al2O3 + 3 H2
leading to the rupturing of Aluminum oxide layer and exposing Al to the action of water:
Al + 3 H2O = Al(OH)3 + 3 H2 (g)
It has now been an hour and the activity has declined but it is much harder to examine the solution owing to a significant presence of a suspensed
salt with the whole solution possessing a metallic Aluminum coloring. Next I added 60 cc of 3% H2O2 and the solution turns white! I heated loosely
covered for 20 seconds again and some activity has returned to the mix. The reaction at this point should follow the Al air battery half reactions:
Anode: Al(s) + 3OH−(aq) → Al(OH)3(s) + 3e−
Cathode: 2 H2O2 + 2H2O(l) + 4e− → 4OH−(aq) + 2 H2O
Overall: 4Al(s) + 6H2O2 → 4Al(OH)3(s)
where I have replaced for O2 our primary oxygen source, H2O2.
I will see if even more of the mix of Al/Al2O3 moves into solution!
[Edited on 24-6-2014 by AJKOER]
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aga
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There is all sorts of stuff in 3% household ammonia.
It's the one liquid [in my experiment] that was the hardest to tell if anything was happening at all, due to the soap bubbles etc.
What are we actually trying to prove here ?
Maybe a better experimental setup needs designing, if there's a point to it.
If not, let's do it anyway !
Anyone seen this : http://www.youtube.com/watch?v=iIq9VvZeWKo
Hydrogen from Al and water proposed as a fuel.
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AJKOER
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One point is that it is an apparent view that weak ammonia , for one, will not rapidly move Al/Al2O3 into solution. There is one source that claims
that weaken gamma alumina containing Al does react at modest (as high as 9) pH. I think some experimental evidence with weak bases (via experiments
with aqueous ammonia and Sodium carbonate) has so suggested such under highly ionic conditions below 100 C indicating something other than a slow
reaction can occur.
Another point is that the electrochemistry of an Aluminum air battery employing H2O2 can also contribute to the rapid dissolution of heated Al/Al2O3.
Crevice corrosion and pitting corrosion associated with air and sea water contact with Aluminum (and some of its alloy) has long been known, but
accepted as slow processes.
[Edited on 24-6-2014 by AJKOER]
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blogfast25
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You can get household ammonia that hasn't been adulterated with detergents, aromas, colourants etc. Easily.
You can also make quite pure NH3 solution up to 20 % NH3 easily yourself: dry distillation of any old ammonium salt with NaOH, KOH or Ca(OH)2.
The 'small problem' with this hydrogen from nanoparticle aluminium spiel is the production of... nanoparticle aluminium. Note that the video doesn't
say one iota about that.
[Edited on 24-6-2014 by blogfast25]
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AJKOER
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Actually, at times, aqueous ammonia with surfactants is better for the task at hand!
I recall reading on the preparation of nano particles via ammonia and it turns out that when surfactants are present, you get smaller particles!
In the current application, as surfactants are mostly composed of common cleaning agents like Sodium stearate, and as there is an issue of the anodes
becoming clogged, I think one would be pleased to have some around to mitigate the reported issue.
However, I am willing to retest with pure aqueous ammonia (which one can prepare from the household brands by placing in a large vessel a plate or
bowl with distilled water and a separate cup filled with the 'impure' ammonia).
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aga
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Seems strong ammonia solution doesn't either (sample 5).
What i saw was pathetically weak gas evolution, almost indistinct from trapped air bubbles until the NaCl was added.
As seen many times before, the oxide gets stripped off rapidly by the Chlorine ions, allowing the Al to react.
Compared to say Al + H2SO4, this reaction is almost a non-event after the NaCl is added.
What were you actually hoping to get as a result from this ?
Aluminium that smells of ammonia ?
Huh ?
They buy it in bulk parcels of 0.1g from the Nanopeople, obviously.
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AJKOER
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One thing people should learn is that Aluminum corrosion can be accelerated both by chemical and electrochemical paths. Mixing metals is asking for
galvanic corrosion.
Here is a near billion dollar mistake paid for by the US Navy on its new Aluminum ship (link: http://mobile.bloomberg.com/news/2011-06-17/navy-finds-aggre... ).
The people responsible for this $$$$$ blunder still apparently think it can be fixed! There will have plently of chances to experiment as this ship is
but one of many in a $30 billion fleet.
[Edited on 24-6-2014 by AJKOER]
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aga
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That's nice dear.
So, er, What was the point of this experiment ?
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AJKOER
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To have fun and dissolve something of course!
In Energetic, to have fun and blow up something of course!
In Chemistry in General, to pretend to have fun and ask why it doesn't work of course!
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aga
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Aha !
Well, i like doing experiments, and did enjoy doing this one.
Dissolve Al/Al2SO3 in HCl or H2SO4 for more fun, and remember to add a pinch of salt 
[Edited on 24-6-2014 by aga]
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