Sciencemadness Discussion Board - Path to Al Acetate and a Tinge of Cyanide(?) via Al Foil/ Carbon/ 5% Vinegar/ 3% H2O2/ Sea Salt and Delayed NH3 (aq) Addition

Path to Al Acetate and a Tinge of Cyanide(?) via Al Foil/ Carbon/ 5% Vinegar/ 3% H2O2/ Sea Salt and Delayed NH3 (aq) Addition

AJKOER - 24-11-2018 at 08:48

This thread’s original intent was to explorer the possible formation of aluminum acetate(s) by the direct action of aluminum metal (sourced from Aluminum foil that has been heat treated to redness by a gas flame) acted upon by household vinegar (5% acetic acid) in the presence of 3% hydrogen peroxide, a graphite rods (sold for use in pencils), and a small amount of sea salt. The mix was heated in a microwave for several 30 second sessions which resulted in a limited formation of a fine white power. Acceleration of the attack of the aluminum metal apparently occurred upon the further addition of pure ammonia water to the mix followed by microwave heating. This is not surprising given that the literature notes the corrosive action of ammonium chloride on aluminum (and other metals, see https://www.researchgate.net/publication/271164226_Effect_of... ).

I believe the underlying reaction to have radical component and can be viewed as a takeoff of my thread ‘Acidic Radicals Based Synthesis’ at http://www.sciencemadness.org/talk/viewthread.php?tid=94166#... .

BACKGROUND
For background I start with Wikipedia comments on Aluminum tri-acetate preparation (link: https://en.wikipedia.org/wiki/Aluminium_triacetate ), to quote:

“According to the CRC Handbook of Inorganic Compounds, aluminium triacetate is a white, water-soluble solid and is usually prepared from aluminium chloride or directly from aluminium by heating in an acetic acid solution with acetic anhydride.[1]
3 CH3COOH + AlCl3 → Al(CH3CO2)3 + 3 HCl
6 CH3COOH + 2 Al → 2 Al(CH3CO2)3 + 3 H2
Theoretically all of the aluminium / acetate / hydroxide salts can be prepared from aluminium hydroxide or sodium aluminate and acetic acid, but formation of the triacetate only occurs in the absence of water.[4] In solutions, the diacetate is the major product formed...”

My presented path could be viewed, in part, as noted above via an in situ created Al(OH)3 (from the action of hydroxyl radicals as derived from H2O2). Suggested path:

Al/Carbon/CH3HCO2-/H2O2/Sea Salt System:
------------------------------------------------------------
2 x [ Al --> Al(lll) + 3 e- ]
6 x [ e- + nH2O --> e-(aq) ]
3 x [ H2O2 = H+ + HO2- ]
3 x [ H+ + e-(aq) = •H ]
3 x [ •H + HO2- = •OH + OH- ]
3 x [ •OH + e-(aq)- = OH- ]
-----------------------------------------------
2 Al + 3 H2O2 ---> 2 Al(OH)3

Then, the action of dilute acetic acid acting on freshly created Al(OH)3.

A confirmation of the above path can be find in a source. To quote from ‘Kinetics of Corrosion Inhibition of Aluminum in Acidic Media by Water-Soluble Natural Polymeric Pectates as Anionic Polyelectrolyte Inhibitors’ by Refat M. Hassan and Ishaq A. Zaafarany, in Materials (Basel). 2013 Jun; 6(6): 2436–2451, doi: 10.3390/ma6062436, link: https://www.ncbi.nlm.nih.gov/pmc/articles/PMC5458939/ :

“The corrosion rate was found to be a function of the concentration of the acid. This result indicates that at least one of the corrosion paths of dissolution of Al metal in HCl solution should involve the presence of hydrogen ions in the rate-determining step…
3.1. Corrosion Mechanism
We propose a suitable mechanism of corrosion, in accordance with the above experimental observations. The corrosion of metal involves an electrochemical process [47,48,49] resulting from dissolution of Al metal in the acid. This process can be expressed by the anodic and cathodic processes, which are defined by Equations (4) and (5), respectively,
Al(s) <--Ox--> Al3+ + 3 e- (4)
2 H+ + 2 e- <--Red--> H2 (5)
The overall electrochemical process can be written as follows:
(6) 2 Al(s) + 6 H+ <-->2 Al3+ + 3 H2(g) (6)
The cathodic reaction produces Hchemisorbed by picking up an electron that released in the anodic reaction (H+ + e = Hchemisorbed ) in Al corrosion in HCl. In such acidic solutions, the Hchemisorbed on the metal surface reacts by combining with other adsorbed Hchemisorbed to form H2 gas molecule, which bubbles from the metal surface. A very small amount of the uncombined Hchemisorbed will remain; however, this amount does not affect the whole process. “

Further possible path explaining the increased corrosion of Aluminum (creating Al ions) in the presence of ammonium ion:

Al = Al(lll) + 3 e-
3 x [ NH4+ = NH3 + H+ ]
3 x [ e- + H+ = •H ]
3 x [ NH3 + •H (or •OH) = •NH2 + H2 (or H2O) ]
Net: Al + 3 NH4+ = Al(lll) + 3 •NH2 + 3 H2 (g)

With the creation of the amidogen radical, a possible REDOX with a metal M:

•NH2 + M(ll) + H+ = NH3 + M(lll) (See, for example, (6) on p. 111O for a Nickel complex at https://www.nist.gov/sites/default/files/documents/srd/jpcrd... )

Interestingly, here is a long article noting "Amine radical cations are highly useful reactive intermediates in amine synthesis" in an organic chemistry setting, see https://www.ncbi.nlm.nih.gov/pmc/articles/PMC3817571/ .

Now, in the case of elemental Aluminum, a corresponding suggested reaction per above would be:

.NH2 + ⅓ Al + H+ = NH3 + ⅓ Al(lll)
(As ⅓ Al --> ⅓ Al(lll) + e-
H+ + e- = •H
•H + •NH2 = NH3 )

In the presence of H2O2, a source of O2 on decomposition, could further interact with any formed amidogen radical as follows:

•NH2 + O2 = •H + HONO

Source: ‘Reaction between the amidogen radical, NH2, and molecular oxygen in low-temperature matrixes’, by John N. Crowley and John R. Sodeau, in J. Phys. Chem., 1989, 93 (12), pp 4785–4790, DOI: 10.1021/j100349a021, link: https://pubs.acs.org/doi/10.1021/j100349a021 or http://www.diva-portal.se/smash/get/diva2:730237/FULLTEXT01.... Table 2.1 on page 25 suggesting products of NO + H2O and in the presence of H2O2 leads also to NO2 in water.

Any formed HONO (or HNO2) is unstable (see https://en.wikipedia.org/wiki/Nitrous_acid ) likely decomposing as follows:

2 HNO2 = NO + NO2 + H2O

As such any NO2 together or HNO2 (especially in the presence of NaCl acting in a similar fashion to HNO3), are reported to attack Aluminum metal. This is likely due to a Fenton-type REDOX reaction in the presence of elemental Aluminum with any created HNO3 or HNO2) See, for example, "Fenton chemistry in biology and medicine*" by Josef Prousek, to quote reaction (15) on page 2330:

"For Fe(II) and Cu(I), this situation can be generally depicted as follows [20,39],

Fe2+/Cu+ + HOX → Fe3+/Cu2+ + •OH + X- (15)

where X = Cl, ONO, and SCN. "

And my prior comments at http://www.sciencemadness.org/talk/viewthread.php?tid=30473#... .
Less likely but possible another reaction of interest includes creation of the acetate radical anion (see https://pubs.acs.org/doi/abs/10.1021/ja00105a032?journalCode... ):

CH3CO2H <-> H+ + CH3CO2-
CH3CO2- + •H --> •CH2CO2- + H2 or •CH2CO- + H2O
•CH2CO2- + •NH2 = CH2NH2- + CO2
Or: •CH2CO- + •NH2 = CH2CONH2

With respect to the last compound, note a recent comment I may on SM referencing cyanide paths (see http://www.sciencemadness.org/talk/viewthread.php?tid=103289 ):

Quote: Originally posted by AJKOER

.....
........
I did notice this article (see 'Excision of CN− and OCN− from acetamide and some amide derivatives triggered by low energy electrons' by Constanze Koenig-Lehmann, et al, abstract at https://pubs.rsc.org/en/Content/ArticleLanding/2008/CP/b8121... ). To quote in part:

"Low energy electron attachment to acetamide and some of its derivatives shows unique features in that the unimolecular reactions of the transient anions are remarkably complex, involving multiple bond cleavages and the formation of new molecules. Each of the three compounds acetamide (CH3C(O)NH2), glycolamide (CH2OHC(O)NH2) and cyanoacetamide (CH2CNC(O)NH2) shows a pronounced resonance located near 2 eV and decomposing into CN− along a concerted reaction forming a neutral H2O molecule and the corresponding radical (methyl and methoxy). "

Per the above, I would express the basis of a possible reaction as:

Al --> Al(lll) + 3 e-

3 x [ CH3C(O)NH2 + e- --> •CH3 + CN- + H2O ]

Net: Al + 3 CH3C(O)NH2 --> 3 •CH3 + Al(lll) + 3 CN- + 3 H2O
......
[Edited on 10-11-2018 by AJKOER][/rquote]

where the corresponding net reaction in the current case could be:

Net: Al + 3 CH2CONH2 --?--> 3 •CH2 + Al(lll) + 3 CN- + 3 H2O

indicating a possible, but small, cyanide presence.

In the current experiment, the action of the hydroxyl radical on elemental carbon is a path to CO:

•OH + C = •H + CO (EDIT see comments and references at http://www.sciencemadness.org/talk/viewthread.php?tid=97845#... )

Also, given the possible presence of CO, from the same referenced SM thread above:

[Edited on 4-10-2014 by AJKOER][/rquote]
--------------------------------------------------------------------

A bit harder, try this path (see https://chemiday.com/en/reaction/3-1-0-261):

CO + NH3 --500 C, Al2O3--> HCN + H2O

"Carbon monoxide react with ammonia to produce hydrogen cyanide and water. The technical method production hydrogen cyanide. This reaction takes place at a temperature of 500-800°C, an overpressure. In this reaction, the catalyst is can be V2O5, CeO2 Al2O3, ThO2."

Text sounds like a translation.

At such high temperatures, I would expect:

NH3 + Heat ---> •H + •NH2

based on the action of hv on ammonia (see R1 at https://journals.ametsoc.org/doi/pdf/10.1175/1520-0469%28197... )

Then, subsequent reactions with CO forming HCN and H2O.

------------------------------------------------------------------

The above speculated radical mechanism is interesting, if correct, as the hydrogen atom radical can be formed at RT by the action of NaOH (or HCl) on Aluminum metal where some .H radical could be imbued on the surface of the Aluminum. My prior related comment:

Quote: Originally posted by AJKOER

.......
Next, imbue the surface of Mg or Al with the hydrogen atom radical (from the traditional nascent hydrogen generation methods based on say Al/NaOH).

One may assume that the •H radical functional behaves (per its seemingly reversible formation reaction: e- + H+ = •H ) as apparently a (e-,H+) pair acting on ions. For an example from 'Hydrometallurgy 2008: Proceedings of the Sixth International Symposium', p. 818, a commercial reductive leaching equation, to quote:

" PbS + 2 •H = Pb + H2S (5) " (see https://books.google.com/books?id=1etfSdk55SYC&pg=PA818&... )

which I view functionally as follows:

Pb(+2)S(2-) + 2 (e-, H+) = Pb + H2S (g)
......
[Edited on 4-10-2018 by AJKOER]

So the surface .H could be further enlisted in other (unspecified) reactions with CO per above via the radical reaction:

•H + NH3 = H2 + •NH2

which are the same radicals (.H, .NH2) postulated above to be formed at high temperatures that eventually result in a HCN product.

[Edited on 9-11-2018 by AJKOER]

Some possible paths to a cyanide presence. Assuming the presence of CO, I would suggest also the simple path:

CO + •NH2 = •CONH2 (or •CN + H2O)
•H2N + CO = •H2NCO (or •NC + H2O)

•H2NCO + •CONH2 = H2NC(O)C(O)NH2 (oxamide)
(Or: •NC + •CN = NCCN )

Per Wiki on Cyanogen (https://en.wikipedia.org/wiki/Cyanogen ) to quote:

“Cyanogen is the anhydride of oxamide:
H2NC(O)C(O)NH2 → NCCN + 2 H2O
…….
Like other cyanides, cyanogen is very toxic, as it readily undergoes reduction to cyanide, …”

Given the theory presented above, the observed pale blue coloration developed in the synthesis following the introduction of pure aqueous NH3 in distilled water, may be due to HCN or a harmless bluish tinge created from the addition of aqueous ammonia to an aluminum salt (see https://mysite.du.edu/~jcalvert/phys/alumin.htm and search on 'bluish').

In any event, perform the experiment in a well vented area.

I believe the reduction in color upon boiling with added vinegar, suggests a successful preparation with a minor threat of HCN.

EXPERIMENT DETAILS
----------------------------------
Started with 200 mL of 5% white distilled household vinegar, 100 mL H2O2, ¼ and increased to ¾ teaspoon of over the counter sea salt, and flame heated Al foil with a preheated surface area of 10.7 sq inches. Also note, I was short on my supply of carbon and employed just 2 small low surface area graphite rods. Ideally, one would like to have a ratio of Aluminum surface area to C that is small to encourage corrosion chemistry with the Al, and not otherwise to foster reactions leading to, say CO, at the more noble carbon electrode. Also, activated carbon (AC) may be superior to graphite as per this source (https://www.researchgate.net/publication/239705822_Generatio...) hydroxyl radicals are created in aqueous solutions with microwave radiation treatment.

[Edited on 24-11-2018 by AJKOER]

[Edited on 25-11-2018 by AJKOER]

AJKOER - 24-11-2018 at 08:54

Start and near finish pictures.
Note, Aluminum consumed except for small piece that was not heat treated.

Photo Editor-20181124_104930.jpg - 449kB

Photo Editor-20181124_104959.jpg - 352kB

[Edited on 24-11-2018 by AJKOER]

AJKOER - 24-11-2018 at 09:05

Final picture taken in a larger vessel (due to the added vinegar) taken some 12 hours after decanting, adding more vinegar and heating. Note, blue marks in the picture are actually on the exterior of the vessel.

Photo Editor-20181124_120119.jpg - 246kB

[Edited on 25-11-2018 by AJKOER]

AJKOER - 24-11-2018 at 09:18

More earlier pictures

Photo Editor-20181124_105046.jpg - 351kB

UC235 - 24-11-2018 at 09:38

What is this incoherent garbage. You corroded some aluminum foil in salty vinegar and made aluminum hydroxide suspension and a mess.

unionised - 24-11-2018 at 09:41

"aluminum metal (sourced from Aluminum foil that has been heat treated to redness by a gas flame)"
Why bother?
Mind you, that question can be asked of the whole thread.

AJKOER - 24-11-2018 at 09:56

Quote: Originally posted by unionised

"aluminum metal (sourced from Aluminum foil that has been heat treated to redness by a gas flame)"
Why bother?
Mind you, that question can be asked of the whole thread.

As to why aluminum foil should be heated in a flame (in places till it is red), apparently, the heat treatment weakens the protective coating and disrupts the annealing process (which makes the Al foil seemingly inert). Here is an extract of my prior comment on this topic:

Quote: Originally posted by AJKOER

.....
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]

The above coupled with several experimental verifications on my part confirm the theory.

But of course, you can easily verify it yourself as either the Al foil rapidly dissolves in some situations or it does not!

[Edited on 24-11-2018 by AJKOER]

AJKOER - 24-11-2018 at 10:40

Note, the path I outlined appears (as all the Al is gone!) effective and does not generate a large amount of the insoluble aluminum diacetate as mentioned in this thread https://www.reddit.com/r/chemistry/comments/4m6to8/make_alum... .To quote:

"If Al(CH*3CO2)3* (aluminium tri-acetate) is what you need, then trying to make this from aluminium metal and acetic acid is not the easiest or most efficient method. It is better to react the aluminium metal with a solution of sodium hydroxide first to make a solution of NaAl(OH)4 (sodium aluminate). This way, the aluminium oxide is not a problem, since it reacts with NaOH solution to make the same product:

Al*2O3* + 2 NaOH + 3 H*2O -----> 2 NaAl(OH)4*

This is quite slow and doesn't produce any gas, so it can look like the reaction is doing nothing while the aluminium oxide dissolves. But once the oxide is all gone, the aluminium metal underneath reacts to form more sodium aluminate. This reaction is much faster:

2 Al + 6 H*2O + 2 NaOH -----> 2 NaAl(OH)4* + 3 H*2*

Once you have a solution of sodium aluminate, the next step is to add 3 molar equivalents of acetic acid:

NaAl(OH)4 + 3 CH*3CO2H -----> CH3CO2Na + Al(CH3CO2)2*OH

So why not add 4 equivalents of acetic acid to make Al(CH*3CO2)3? Because aluminium tri-acetate is water soluble, while aluminium di-acetate is not. So in the reaction above, the Al(CH3CO2)2OH precipitates as a solid, while the CH3CO2*Na (sodium acetate) byproduct stays in solution.

So you can then isolate pure Al(CH*3CO2)2*OH and then add this to one more equivalent of acetic acid to make aluminium tri-acetate solution:

Al(CH*3CO2)2OH + CH3CO2H -----> Al(CH3CO2)3* + H*2*O

Then to make solid Al(CH*3CO2)3*, just leave the solution to evaporate (if you want nice crystals) or boil the water off to make a powder."

[Edited on 24-11-2018 by AJKOER]

AJKOER - 24-11-2018 at 10:58

Quote: Originally posted by UC235

What is this incoherent garbage. You corroded some aluminum foil in salty vinegar and made aluminum hydroxide suspension and a mess.

Thanks for the comment sharing concerns some may have.

With respect to 'incoherent', I refer you to my prior referenced thread employing acidic radicals (link: http://www.sciencemadness.org/talk/viewthread.php?tid=94166#... where interestingly, with the source article of that thread having a web publication date of September 15, 2017, perhaps many textbooks are so outdated, they could be 'garbage').

Thanks also in mentioning the presence of salt, in particular sea salt, which apparently has many roles. First, as my prep is based on an electrochemical cell (or battery cell), it is in the province of electrochemistry and can be referred to as an electrosynthesis (see https://en.wikipedia.org/wiki/Electrosynthesis). In that regard, NaCl serves in the role of an electrolyte. Second, in the corrosion chemistry of Aluminum, NaCl can apparently engage in surface attacks of the protective Al2O3 layer (see this review for more details http://jes.ecsdl.org/content/161/9/C421.full ). Lastly, not pure NaCl, but sea salt, which some researchers have noted interestingly that it may be chemically more active (perhaps from multiple trace elements which can be catalytic and I would support this conjecture by the apparent biochemical sensitivity of fish to replacing sea salt with pure NaCl, read https://books.google.com/books?id=H1BGAAAAYAAJ&pg=PA72&a... ).

Now, the comment may be partially correct on the aluminum hydroxide suspension perhaps created upon the addition of aqueous NH3, due to the bluish tint (and perhaps not the more problematic HCN), per this source (https://mysite.du.edu/~jcalvert/phys/alumin.htm ) to quote:

"Aluminium dissolves slowly in dilute hydrochloric acid to make a clear solution. When ammonium hydroxide is added, a characteristic translucent gel precipitates, with a bluish tinge."

However, upon decanting to remove most of the mess, and then the reduction of color upon addition of more acetic acid with microwave heating, implies to me likely removal of the Al(OH)3 translucent gel or other created compound.
-------------------------------------------------

Anyone, however, is free to follow the CRC Handbook of Inorganic Compounds, to prepare aluminium triacetate from aluminium chloride or directly from aluminium by heating in an acetic acid solution with acetic anhydride.

[Edited on 25-11-2018 by AJKOER]

AJKOER - 25-11-2018 at 04:58

Both pictures taken 1 day after decanting. Note, a small amount of fine white precipitate (possibly Aluminum diacetate) and the bottom picture (taken in sunlight minutes latter) displays a green tint (ferrous (?) from small iron content of the Al foil or found in vinegar). Solution shows not sign of viscosity.

Photo Editor-20181125_075326.jpg - 288kB

Photo Editor-20181125_075251.jpg - 197kB

[Edited on 25-11-2018 by AJKOER]

unionised - 25-11-2018 at 05:52

Quote: Originally posted by AJKOER

The process for annealing aluminum is to heat it near to the melting point.
So, your idea of "disrupting" the annealing, by annealing it makes no sense.

AJKOER - 25-11-2018 at 07:10

Quote: Originally posted by unionised

.....
The process for annealing aluminum is to heat it near to the melting point.
So, your idea of "disrupting" the annealing, by annealing it makes no sense.

Yes, but a better path to growing a strong protective Al2O3 layer is via anodizing (see https://www.thebalance.com/what-is-anodizing-2340009 ) and not heating.

Also, the qualification is heating till the Al foil glows red (high temperature induced disruption I would suspect).
---------------------------------------------------------

A corrosion related point, I claim my ammonia water was likely pure NH3 (aq) as I prepared it by placing distilled water (DW) in a large mouth vessel together in a container (which was twice the volume of the DW) loaded with yellow colored household ammonia rich in surfactants, etc. I carefully placed a plug (prepared by wrapping layers of Al foil around my finger) full of NaCl into the household ammonia and sealed the entire chamber. As NH3 has much reduced solubility in seawater, I expected with time the DW would amass a good amount of the liberated ammonia from the now salt rich household ammonia.

A day latter, I opened the vessel, I was surprised to see the amount of bubbling/attack performed on the Al plug! This level of corrosion by NH3/NaCl is why I added the pure NH3 water to this threads reaction mix as I was not impressed by the level of attack I was seeing. The result is apparent!

As to the mechanics of the attack by NH4+, one source notes 'a pH and corrosion potential changes with time', see http://citeseerx.ist.psu.edu/viewdoc/download?doi=10.1.1.849... . My suggested radical attack either directly by the •NH2 radical, or from products generated from associated REDOX reactions, is my assessment, which is in line with this observation.

Further support for my radical based path is a case from 1983 (see https://www.researchgate.net/publication/271683294_Stabiliza... ) of stabilizing corrosion in the presence of both Al and Cu. The answer is washing with an aerated mix of NH3/NH4+. My take on the mechanics is that the Cu/NH3/O2 battery produces solvated electrons, which in the presence of O2, makes superoxide (•O2-). This latter radical anion is great in scavenging many radicals (including in the human body) as illustrated as follows:

•NH2 + •O2- + H+ (from water) = NH3 + O2

•OH + •O2- = OH- + O2

•CO3- + •O2- = CO3(2-) + O2

Also, a reaction with Al3+, on the Al surface, which could be impeding (by creating the aluminum semi reduced superoxide product below, which is capable of converting metal ions to low valence states, and associated levels of solubility like with cuprous salts, see cited reference below):

Al3+ + •O2- = [Al(•O2-)](2+) (see comments and sources at http://www.sciencemadness.org/talk/viewthread.php?tid=96347#... )

all of which may reduce the progress of a corrosive attack.

[Edited on 25-11-2018 by AJKOER]

DraconicAcid - 25-11-2018 at 10:47

Quote: Originally posted by unionised

So, your idea of "disrupting" the annealing, by annealing it makes no sense.

Why would it be different from any of his other "ideas"?

Tsjerk - 25-11-2018 at 12:15

That, and there is no protective coating on aluminium foil, except for the couple of nano meters of aluminium oxide. Heating it over a flame will make it brittle and effectively convert it to aluminium oxide. Try to measure volume of hydrogen gas evolved with NaOH with and without heating it to red dullness if you don't believe me.

AJKOER - 25-11-2018 at 12:31

Quote: Originally posted by DraconicAcid

Quote: Originally posted by unionised

So, your idea of "disrupting" the annealing, by annealing it makes no sense.

Why would it be different from any of his other "ideas"?

Thanks Draconic acid but here is an extract of the original discussion as to the process:

Quote: Originally posted by AJKOER

.....
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:

"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.

[Edited on 24-11-2018 by AJKOER]

So, there is a white paper quoted again above referring to a mechanochemical process that starts on page 8.

Then, on page 9, there is then a diagram in the white paper from which I reprinted the author's cited reaction path leading to H2 and bursting of the Al2O3 layer.

Sorry, but I don't see my 'idea', and I doubt anyone wants to assess my artistic talent in attempting to reproduce that figure.

The 2008 work I am quoting from is titled 'Reaction of Aluminum with Water to Produce Hydrogen' by John Petrovic (Ph.D, Materials Science, see https://www.linkedin.com/in/john-petrovic-49aa8241 ) and George J Thomas, who are consultants to the U.S. Department of Energy on the Hydrogen Program. Both authors are retired scientists from Los Alamos National Laboratory and Sandia National Laboratories, respectively.

[Edited on 25-11-2018 by AJKOER]

Texium - 25-11-2018 at 12:55

Alright. Let's take some time to pick this one apart.

Quote: Originally posted by AJKOER

"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."

The only thing the authors describe here in detail that relates to heating is heating a mixture of aluminum powder and a particular type of fine aluminum oxide powder. This is an entirely different situation than sticking a piece of aluminum foil in a flame. Without any actual analysis of your supposed "flame activated aluminum" and more conclusive results, the paper that you cite has absolutely NO relevancy to what you're talking about.

Quote: Originally posted by AJKOER

The problem is that you extrapolated from an experiment that is entirely different than what you conducted, and tried to use that as evidence for your dubious and unquantitative claims, and then you took the assumptions that you made from that and made more assumptions that you attempt to back up with more papers that would only be even tangentially related if some of your assumptions were correct (which they aren't).

You are wrong on so many levels that this barely scratches the surface. Maybe you have some false impression that the crackpot experiments and conjecture that you often post here go unchallenged because everyone believes you, or people don't know how to refute you. The truth is though, it's just exhausting, and nobody wants to waste their time doing it when you're just going to go on believing your fantasy version of chemistry and never actually learn a damn thing.

Tsjerk - 25-11-2018 at 13:10

Quote: Originally posted by Texium (zts16)

What you describe is his hallmark indeed.

Quote: Originally posted by Texium (zts16)

You are wrong on so many levels that this barely scratches the surface. Maybe you have some false impression that the crackpot experiments and conjecture that you often post here go unchallenged because everyone believes you, or people don't know how to refute you. The truth is though, it's just exhausting, and nobody wants to waste their time doing it when you're just going to go on believing your fantasy version of chemistry and never actually learn a damn thing.

I tried to refute him, but he keeps, and keeps, and keeps coming with crap about why he isn't wrong. I can't imaging how much time AJOEKER needed to write his 2400+ posts... Normally you would think someone with a name like AJOEKER is "a joker", but that would be one hell of an elaborate "joke".

I didn't know there was a term like fractal wrongness, I love it and it spot on in this case.

Quote:

Fractally wrong people are often immune to the stopped clock rule because they are not exactly stopped clocks. More like clocks losing a random number of seconds a day, in the wrong time zone of the wrong planet.

LMAO

[Edited on 25-11-2018 by Tsjerk]

AJKOER - 25-11-2018 at 13:51

Your comments are technically accurate, however, I have used and learned to use flame heated aluminum by experimenting. Note, Al foil can also have a plastic (acrylic) coating at times (discussed in an old thread on SM here http://www.sciencemadness.org/talk/viewthread.php?tid=72608#... ), so flame treating also removes any other protective coating.

The purpose for heat treating the Al foil becomes especially obvious with the so called Bleach Battery. Mix aqueous NaOCl with a weak acid (like vinegar) forming HOCl. Pour it into two vessels, each with copper pennies and one with Al foil that is heat treated and the other that is not treated.

It becomes quickly apparent that there is the significant different in chlorine formation, so perform this experiment in a well ventilated setting.

Note, one half cell reaction involves Al, the other HOCl (which breaks down into Cl2). Believe me there is an obvious difference in the speed of the reaction in the vessel with heat treated Al foil in releasing chlorine fumes.

Some comments from a prior thread:

Quote: Originally posted by AJKOER

.......
Note, as I previously performed on SM, chlorine gas can be sourced from a 'bleach battery' without the use of a strong acid. Here is a detailed rehash of the steps involved in the so called Bleach battery:

1. Whip up some Hypochlorous acid by mixing bleach (NaOCl) and vinegar (which contains Acetic acid HAc) in the volume ratio 1.4 parts of 5% vinegar to one part of 8.25% extra strength chlorine bleach.

2. To the HOCl add a piece of copper metal which will function as the cathode and an Aluminum source (like foil, but heat to red on a stove to increase reactivity) to act as the anode.

3. Lastly, add much NaCl to act as the electrolyte and to salt out the chlorine. I usually jump start the reaction in a microwave (1 minute should work).

The underlying chemical reaction leading to hypochlorous acid is given by:

NaOCl + HAc --> HOCl + NaAC

-------------

An alternate preparation of HOCl creating more conc HOCl (which is also avoids organic acids subject to attack by the hypochlorous acid in the presence of copper) is:

2 NaOCl + CaCl2 = 2 NaCl + Ca(OCl)2
Ca(OCl)2 + 2 NaHCO3 --> Na2CO3 + CaCO3 (s) + 2 HOCl (cool and let stand to remove the CaCO3)

Net: 2 NaOCl + CaCl2 + 2 NaHCO3 --> 2 NaCl + Na2CO3 + CaCO3 (s) + 2 HOCl
------------------------

My take on the major electrochemical half-reactions:

In anodic zone (aluminum pieces):

6 H2O <--> 3 H3O+ + 3 OH-

Al + 3OH- ⇒ Al(OH)3 + 3 e-

At the cathode (copper metal):

3 HOCl + 3 H3O+ + 3 e- ⇒ 3/2 Cl2(g) + 6 H2O

for an implied net reaction of:

3 HOCl (aq) + Al (s) --NaCl--> Al(OH)3 (s) + 3/2 Cl2 (g) Eo net = 3.93 V

Note, the moles of consumed HOCl (from NaOCl) to form one mole is not as efficient as employing, for example, NaHSO4 acting on NaOCl+NaCl, but no strong acid or acid salt was used!

Incidentally, the battery cell is theoretically capable of generating 3.93 volts. References: see http://www.exo.net/~pauld/saltwater/ and http://sci-toys.com/scitoys/scitoys/echem/batteries/batterie... and also http://www.dtic.mil/dtic/tr/fulltext/u2/d019917.pdf

[Edited on 9-7-2018 by AJKOER]

Tsjerk - 25-11-2018 at 14:21

You are a caricature of yourself.

AJKOER - 25-11-2018 at 15:42

Quote: Originally posted by Tsjerk

You are a caricature of yourself.

I have since added the credentials of one of the authors, whose work was assumed to be mine (since I did not place quotes around my transcription of his reactions depicted in a diagram) and now responsible for my ill fame, relating to mechanochemical treatment of aluminum metal at elevated temperatures. He is John Petrovic, Ph.D, Materials Science, see https://www.linkedin.com/in/john-petrovic-49aa8241 .
------------------------------------------------------------------

Tsjerk, I wonder if you have the ability to do even the simple experiment I detailed (don't forget the 1 minute microwave jump start), which should take no more than 10 minutes to get it all together. If so you would be still smelling the chlorine (from the chamber with the treated Al foil)!
-----------------------------------------

A bigger challenge, repeat my thread synthesis replacing the graphite with a lot of activated carbon (AC), so that its surface area is in excess of the Aluminum, employing a microwave and don't forget the sea salt!

I would definitely expect even better results, based on this 2007 work (has anyone noticed how recent the science is I am citing?, link: https://www.researchgate.net/publication/239705822_Generatio...

My take:

AC = AC+ + e- (aq)

Then, as in the opening part of this thread:

Al/Carbon/CH3HCO2-/H2O2/Sea Salt System:
------------------------------------------------------------
2 x [ Al --> Al(lll) + 3 e- ]

6 x [ e- + nH2O --> e-(aq) ]
3 x [ H2O2 = H+ + HO2- ]
3 x [ H+ + e-(aq) = •H ]
3 x [ •H + HO2- = •OH + OH- ]

Memo: Net last 4 reactions:
e-(aq) + H2O2 = •OH + OH- (a one electron reduction of H2O2 likely accelerated by microwaving in the presence of AC)

3 x [ •OH + e-(aq)- = OH- ]
----------------------------------------------------
2 Al + 3 H2O2 + nH2O ---> 2 [Al(H2O)6](OH)3

Then, the action of dilute acetic acid acting on freshly created Aluminum hydroxide aqua complex, plus likely a needed ammonia/NH4+ boost via:

Al = Al(lll) + 3 e-
3 x [ NH4+ = NH3 + H+ ]
3 x [ e- + H+ = •H ]
3 x [ NH3 + •H (or •OH) = •NH2 + H2 (or H2O) ]
-------------------------------------------------------
Net: Al + 3 NH4+ = Al(lll) + 3 •NH2 + 3 H2 (g)

Not really that complex (at least, so far).

[Edited on 26-11-2018 by AJKOER]

woelen - 26-11-2018 at 00:15

I tried to understand the point of thios experiment, but I lost track of it almost immediately at the start of the first post

@AJKOER: In nearly all of your posts you complicate things so much that it becomes gibberish for most (all?) readers of this forum. Who can assess the correctness of what you write? Who wants to wade through tons and tons of seemingly irrelevant equations and references, often presented in such a way that there hardly is any visible structure in the "dumps".

If you want people to show any interest in your experiments or discussions, then you really have to change things! Most people simply ignore your posts and unfortunately they have a good reason to do so. You are the one who can change that, only you!

AJKOER - 26-11-2018 at 02:10

Thanks Woelen, I will separate out the highly technical references for those interested.
--------------------------------------------------------------------------------------------------

Back to my challenged claim that heated Al foil can produce a gamma alumina coating that has been described by an authority (John Petrovic, Ph.D, Materials Science) as having a weakened surface structure, and further per the web:

"The diffraction patterns in the temperature range (550–850 °C) correspond to the γ-Al2O3 phase."

Also, per the web:

"Aluminum melts at around 600 C and steel around 1400 C. Molten steel is white hot and molten aluminum red hot. "

So, red hot Al is in the range of possible gamma alumina creation, which may induce a weakened gamma layer.

In my opinion, several prior experiments performed by me confirm this reactivity as well, and if anyone wants to continue to challenge the reactivity claim, perhaps try to perform my suggested experiment relating to a Bleach Battery cell. However, one need only look at the pictures in this very thread and my prior comment, to quote from caption on one of the thread pictures:

"Note, Aluminum consumed except for small piece that was not heat treated."

Apparently, I was not able to uniformly heat all the Al foil to redness leaving a small untreated piece that is still resistant.

However, if one still wishes to contest this point, it does not matter to the central fact that in this thread, using microwave (MV) assisted heating (perhaps key per a reference claiming radical formation with MV application on activated carbon and other sources citing interaction between hydroxyl radicals and graphite), I fairly rapidly dissolved heat treated Al foil (using a methane flame) with weak 5% vinegar, weak 3% H2O2, two small graphite rods and sea salt followed by a late addition of dilute pure ammonia water, likely producing not Al(OH)3, but an aluminum acetate salt (to be determined).

I added some alcohol to a sample of my final mix and eventually it should produce crystals. Unfortunately, as I increased NaCl dosing early on (to try to dissolve the Al foil before I conceived of adding NH3), I expect lots of NaCl, but in future runs, this should not be an issue.

[Edited on 26-11-2018 by AJKOER]

Tsjerk - 26-11-2018 at 08:34

Quote: Originally posted by AJKOER

Tsjerk, I wonder if you have the ability to do even the simple experiment I detailed (don't forget the 1 minute microwave jump start), which should take no more than 10 minutes to get it all together. If so you would be still smelling the chlorine (from the chamber with the treated Al foil)!

[Edited on 26-11-2018 by AJKOER]

You are suggesting I should microwave a mixture of vinegar, concentrated bleach, salt, copper and burned aluminium foil and then smell it... Thank you very much but I will pass on that one.

I do have some 99% NaOH draincleaner and aluminum foil in my kitchen, I will mix some burned and unburned alu-foil with some dissolved NaOH and report what happens... As my whole point was that aluminium foil in a flame oxidizes to the oxide. It won't produce bubbles anymore as there is no aluminium anymore. You have oxide in your experiments, not aluminium.

AJKOER - 26-11-2018 at 09:00

Quote: Originally posted by Tsjerk

...
....
You are suggesting I should microwave a mixture of vinegar, concentrated bleach, salt, copper and burned aluminium foil and then smell it... Thank you very much but I will pass on that one.

.....As my whole point was that aluminium foil in a flame oxidizes to the oxide. It won't produce bubbles anymore as there is no aluminium anymore. You have oxide in your experiments, not aluminium.

On the Bleach Battery, you can use the microwave for say 40 seconds on both vessels at once. This reduces the inception period, and then see the difference in Cl2 creation. The electrochemical cell does not generate chlorine from Al2O3, so by your argument (lots of Al2O3 created), there should be no contest, and your untreated Al foil should prove to be superior!

My view is it that it not the quantity of Al presence after flame treatment, but the amount of exposed Al feeding the half cell reaction.

Now, per the first picture in this thread, it still looks pretty good. But even if it all Al2O3, should not the oxide be more easily subject to an acid attack creating acetate? I don't agree, but what is the problem if it does dissolve faster, you should endorse the flame treatment process.

[Edited on 26-11-2018 by AJKOER]

Tsjerk - 26-11-2018 at 12:36

Did you do this chlorine battery experiment you talk about with a proper negative?

I just put some "treated" aluminum in NaOH, with a proper negative, and found the burned aluminium to be about 95% reduced in hydrogen production.

walruslover69 - 26-11-2018 at 12:56

Just from skimming through the wiki articles on aluminum acetates, It states that the triacetate hydrolyzes in solution to the mono and di acetate. You could prevent this by using really concentrated acetic acid solution or acetic anhydride, but then you are just producing the triacetate directly without the need for any of the method.

AJKOER - 26-11-2018 at 17:54

Thanks for all the comments to date.

I am working on a new related experiment, easier to read and comprehend, to be released soon which provides more insight into the kinetics!

My alcohol enriched final mix is losing some volume, but it may still take a while to view any crystals.

Interestingly, I was cleaning up a vessel that I left with some residual solution which included the graphite electrode (as part of the decanting to remove the 'mess') and was hit by a strong smell from apparently an electron/radical induced reaction involving acetate and ammonia (strong over powering and a bit sweet smell). I suspect this accidental composition is best avoided.

[Edited on 27-11-2018 by AJKOER]

woelen - 27-11-2018 at 00:36

Quote: Originally posted by AJKOER

[...]Interestingly, I was cleaning up a vessel that I left with some residual solution which included the graphite electrode (as part of the decanting to remove the 'mess') and was hit by a strong smell from apparently an electron/radical induced reaction involving acetate and ammonia (strong over powering and a bit sweet smell). I suspect this accidental composition is best avoided.

Sure that the smell comes from an electron/radical induced reaction? Isn't it simply an impurity in the vinegar used, or in the ammonia, or in the alcohol? What grade chemicals did you use?

Impurities in household chemicals can be anything and they can produce all kinds of smells when allowed to interact with other chemicals for a longer time. No useful data can be derived from such experiments, especially more subtle side effects usually tell nothing about the original reaction, they just come from the impurities.

I have an example of this from my own experience.

One day, I made nickel nitrate by dissolving a few dutch coins (dubbeltje, 10 cents, appr. 99.5% Ni, 0.5% Sn according to info from some numismatic website).
I precipitated the nickel(II) with sodium hydroxide and obtained a green precipitate of Ni(OH)2. Next, I filtered this and added a solution of Na2S2O8. This leads to formation of a black precipitate of some higher oxidation state of nickel (oxidation state +3, maybe even a little +4). Next, I added the black precipitate to an acid solution (dilute H2SO4). And what happens? I get formation of O2, the precipitate dissolves and I get a pink/purple solution. Not a pure pink color, but a somewhat dull/dirty looking pink. I was really stunned about this. How could nickel produce pink solutions?
I soon found out, however, that the pink color was due to manganese impurity and with the strongly oxidizing Na2S2O8 a tiny amount of permanganate ion was formed. Apparently those dubbeltjes contained a tiny amount of manganese, not mentioned in the numismatic websites. They only list what is added in the alloy, but the used metals may contain tiny amounts of impurities. Even 0.01% of manganese will lead to visible pink color when converted to permanganate.

My lesson from this experience is to be very cautious when observing secondary effects, which are totally unexpected. Nearly always these effects are due to impurities, albeit not always. It has become a second nature for me to investigate the occurrence of such secondary effects and only very few times there really is an interesting thing and something new is observed.

MrHomeScientist - 27-11-2018 at 06:21

woelen, how did you find out that manganese was the culprit? Did you do an experiment, or find a more detailed analysis report for the metal?

In either case, this should be a good lesson for Joker. When you see something interesting, you shouldn't immediately jump to some arcane "radical" or "HOCl"-related conclusion just because that's your favorite pet theory. "Huh that smells funny. It must be radicals!"
There's no reason to invoke a million obscure research papers and overly complicated mechanisms when "impurities in your household materials" explains things just as well.

Tsjerk - 27-11-2018 at 08:07

As far as I could figure out AJOEKER uses kitchen vinegar as source of vinegar, which in combination with bleach would give a wide variety of chlorinated organic substances which are known to have a very perceivable smell.

DraconicAcid - 27-11-2018 at 10:13

Quote:

Given the theory presented above, the observed pale blue coloration developed in the synthesis following the introduction of pure aqueous NH3 in distilled water, may be due to HCN or a harmless bluish tinge created from the addition of aqueous ammonia to an aluminum salt (see https://mysite.du.edu/~jcalvert/phys/alumin.htm and search on 'bluish').

In any event, perform the experiment in a well vented area.

I believe the reduction in color upon boiling with added vinegar, suggests a successful preparation with a minor threat of HCN

Your only indication that you produced cyanide was that it was blue, and the blue colour went away when you acidified it?

You do realize that cyanide ion isn't blue, right?

Deathunter88 - 27-11-2018 at 13:06

Quote: Originally posted by MrHomeScientist

In either case, this should be a good lesson for Joker. When you see something interesting, you shouldn't immediately jump to some arcane "radical" or "HOCl"-related conclusion just because that's your favorite pet theory. "Huh that smells funny. It must be radicals!"
There's no reason to invoke a million obscure research papers and overly complicated mechanisms when "impurities in your household materials" explains things just as well.

A million times this. Someone needs to teach AJKOER the concept of Occam's razor. On a side note, why has he not been banned yet?

j_sum1 - 27-11-2018 at 14:13

Quote: Originally posted by Deathunter88

On a side note, why has he not been banned yet?

Level heads prevail.
Even this thread has provoked some good discussion related to sound practice and how to draw appropriate conclusions. I think such discussions are valuable and of particular benefit to newbies. I like it that we can have a civil thread where we call out crap for what it is, and discuss what is reasonable. This kind of thing is rare on the internets.

How much of this AKOJER takes on board is really up to him. I don't see that he has done anything banworthy.

MrHomeScientist - 27-11-2018 at 14:17

I think a clever title for him would be appropriate. Something like deltaH's "Dangerous source of unreferenced speculation".

AJKOER - 28-11-2018 at 05:13

Quote:

Per Draconic Acid:

[rquote]
Given the theory presented above, the observed pale blue coloration developed in the synthesis following the introduction of pure aqueous NH3 in distilled water, may be due to HCN or a harmless bluish tinge created from the addition of aqueous ammonia to an aluminum salt (see https://mysite.du.edu/~jcalvert/phys/alumin.htm and search on 'bluish').

In any event, perform the experiment in a well vented area.

I believe the reduction in color upon boiling with added vinegar, suggests a successful preparation with a minor threat of HCN

Your only indication that you produced cyanide was that it was blue, and the blue colour went away when you acidified it?

You do realize that cyanide ion isn't blue, right?

No, DraconicAcid, you miss a significant part of my thread which I personally felt was needed to address a potential safety issue with the experiment. Here it is:

Quote: Originally posted by AJKOER

Apparently my having opine on possible paths to cyanide (some even recently) and my not connecting dots leading to someone personal injury was not a risk I was willing to dismiss entirely whether it be an immediate or delayed hazard.

And, guess what, the presence of electrons did eventual create some unknown breakdown product! What I failed to mention is how I felt after seeming minor exposure to the fumes, I just decided to issue a warning.
----------------------------------------

By the way this is not the first time working with electrochemical cells and organics produced an unexpected breakdown product (like CHCl3), see http://www.sciencemadness.org/talk/viewthread.php?tid=27530 , partial extract below:

Quote: Originally posted by AJKOER

I would like to interject a point on perhaps an overlooked aspect of my bleach battery cell experiment from a new perspective of totally informed hindsight.

Electrosynthesis can occur not only via standard electrolysis, but apparent also via an in situ formed electrochemical cell as well, per my recollection (see as examples http://edepot.wur.nl/385426 and http://pubs.rsc.org/en/content/articlelanding/2010/gc/c0gc00... ).

In my experiment, I did in fact produced a battery cell, and given the added sea salt which may have promoted some solvated electron activity, one could describe the process as an attempted electrosynthesis path to assist in the cleavage formation of chloroform from citric acid/citrate. Possible reactions of interest would be:

HOCl -> HCl + 1/2 O2 (minor decomposition pathway)

O2 + e-(aq) = .O2- (see Table l, p. 14 at http://iopscience.iop.org/article/10.1088/0022-3727/48/42/42... )

Cl2 + .O2- = .Cl2- + O2 (see Reaction 36 at SUPPLEMENTAL SECTION PARTICIPATION OF THE HALIDES IN PHOTOCHEMICAL REACTIONS IN NATURAL WATERS AND TREATED WATERS by Yi Yang and Joseph J. Pignatello)

HOCl + e-(aq) -> OH- + .Cl

.Cl + Cl- = .Cl2- (see https://books.google.com/books?id=mRoJUB5fxRwC&pg=PA321&... )

To test this hypothesis for those interested, if one repeats my experiment WITHOUT the presence of aluminum metal (which I usually source from aluminum foil heated to red hot forming a more reactive gamma aluminum oxide coating rich in surface defects), one may observe lower or no yield of CHCl3 (as there was no added cupric salt) if any electrosynthesis was likely involved.
-----------------------------------------------------------------

Another......

[Edited on 21-5-2018 by AJKOER]

A word of advice, if working in a setting (high temperature combustion, electrolysis, photolysis, electrosynthesis,...) and you have mixtures of elements like carbon, hydrogen and chlorine, expect possible creation of CHCl3, CxCly,...., as potential breakdown products.

Similarly, mix of carbon, hydrogen and nitrogen, HCN is a frequent product (example, combustion of certain plastics) .

Mix of carbon, hydrogen, nitrogen and oxygen, HCN, as detailed in my thread, is still a possibility.

Mix of carbon, chlorine and oxygen, watch out for COCl2.

Mix of H, P, and O, deadly PH3 may still appear as a breakdown product.

.........

[Edited on 28-11-2018 by AJKOER]

[Edited on 28-11-2018 by AJKOER]

AJKOER - 28-11-2018 at 07:58

Quote: Originally posted by Tsjerk

As far as I could figure out AJOEKER uses kitchen vinegar as source of vinegar, which in combination with bleach would give a wide variety of chlorinated organic substances which are known to have a very perceivable smell.

Sorry, but this thread is really about a path to 'Al Acetate via Al Foil/ Graphite/ 5% Vinegar/ 3% H2O2/ Sea Salt and Delayed NH3 (aq) Addition'.

No use of NaOCl or HOCl. However, the 3% H2O2 may have a bit of H3PO4 as a pH stablizer.

As what is in the Distilled White Vinegar (USA), see https://supremevinegar.com/2016/08/22/white-distilled-vinega... .

In China, to quote(see https://www.quora.com/What-does-vinegar-have-in-it ):

"The main components of Vinegar are Acetic Acid (3%--5%),and it also contains various Amino Acids, Organic, Acids, Sugars, Vitamins, Nutritional Ingredients such as Alcohol and Ester."

[Edited on 28-11-2018 by AJKOER]

AJKOER - 28-11-2018 at 08:51

Woelen your impurity point is valid given I am working with household vinegar. This by preference based on cost and availability produced from chemically pure alcohol, as I understand the history of recent preparations of white vinegar, acted upon by bacteria (as mandated in US) which implies assorted organics.

Let me describe the composition of the residual mix from decanting, which led to the unknown compound, as sourced from unreacted Al foil, vinegar, Al3+, NH4+, NaCl, O2/H2O2 and a graphite rod. If I happen to use Reynolds Aluminum Foil (see http://www.sciencemadness.org/talk/viewthread.php?tid=16337 ) add iron to feed a fenton reaction adding hydroxyl radicals. To the extent that the electrochemical cell is likely still functioning, add some solvated electrons and therefrom associated reactive oxygen species (like superoxide, H2O2,...).

A interesting aspect of a system with Fe ions and any O2 in the presence of Al3+, is the ability of the latter to form an intermediate complex with superoxide (or perhydroxyl) to keeping cycling iron between complexes of Fe3+ and Fe2+, feeding a slow fenton (due to low iron concentration and unfavorable pH), forming associated radicals (see http://www.sciencemadness.org/talk/viewthread.php?tid=96347#... ).

I see the potential for assorted radicals decomposing organics, and further interaction of those radicals, forming in time some stable breakdown products.

[Edited on 28-11-2018 by AJKOER]

AJKOER - 29-11-2018 at 16:59

Just came across this in my readings, to quote:

"In addition, while other oxide nanocrystals were ineffective to promote hydrogen generation in tap water, TiO2 nanocrystals (P90, w14 nm in diameter) were found to be highly effective in facilitating the production of hydrogen from the reaction of Al with tap water, comparable to the well-known g-Al2O3."

Article: 'Generation of hydrogen from aluminum and water - Effect of metal oxide nanocrystals and water quality' by Hong-Wen Wang, et al, in International Journal of Hydrogen Energy, Volume 36, Issue 23, November 2011, Pages 15136-15144, Link: https://depts.washington.edu/solgel/documents/pub_docs/journ... .

Well, apparently, comparable to the well-known gamma-Al2O3 in some circles.

Texium - 30-11-2018 at 06:25

That's all well and good, but you never proved that YOU made gamma aluminum oxide to begin with!

Everything you've been doing is just cargo cult science. You do "experiments" and cite copious references, but upon looking for any sort of scientific rigor among your many lengthy posts, it all falls apart.

If your posts were academic journal articles, they'd be all Conclusion with almost no Experimental and absolutely no Analysis, and frankly, that doesn't work.

Edit- Relevant bit from the Wikipedia page I linked:

Quote:

An example of cargo cult science is an experiment that uses another researcher's results in lieu of an experimental control. Since the other researcher's conditions might differ from those of the present experiment in unknown ways, differences in the outcome might have no relation to the independent variable under consideration.

[Edited on 11-30-2018 by Texium (zts16)]

AJKOER - 30-11-2018 at 07:00

Quote: Originally posted by Texium (zts16)

That's all well and good, but you never proved that YOU made gamma aluminum oxide to begin with!

True, I do not have the means to directly inspect the treated Aluminum and state there exists some gamma aluminum oxide.

However, there is some indirect evidence present, I would suggest, due to the property of the gamma aluminum oxide layer given its weaken structure allowing reaction with tap water, to form hydrogen disrupting the protective layer (here tap water implies to me H2O with transition metal and/or bicarbonate presence, like ferrous bicarbonate).

Bottom line, significant dissolving (or, at least, partial) of the treated Al is possibly suggestive/consistent with a weaken gamma aluminum oxide presence, but not conclusive.

Going forward, I will qualify my reference to an attempt of employing high temperature to induce a gamma aluminum oxide presence as being suggestive (or possible) to a limited extent.
-------------------------------------------------------------------

As to the history as I came to become acquainted with the heat treating process, it starts with much experimental frustration on how apparently inert Al foil was behaving (which is not surprising as the Aluminum wrap product industry survives on a perpetuation of the belief that annealed/acrylic coated Al is relatively inert/safe). The mention of plastic coatings in a thread led to a fire test of a strip of foil. The residue was tested in a Bleach battery cell and was evidently more active in reacting in the system HOCl/Cu/NaCl. Then came a document search to understand possibly why (and also to check the box on plausible explanations), which led to a very little known white paper that was not even precisely on topic, mechanical and high temperature induced layer of gamma alumina on fine Aluminum powders. Research on gamma alumina mentions temperature ranges that are apparently consistent with heating Al to red hot.

No cargo cult science here, as all related experimenting was performed by myself. Even my cited white paper can't be borrowed science as it is in a different specialized field working with powders. I am only suggesting, by itself, a temperature treatment path, which per one posited explanation, ascribes an induced gamma alumina layer on larger pieces of Aluminum. In any account, the process appears to assist in the attack of Al in repeated runs and varying experimental settings (some of which are posted on SM).

[EDIT] Those assuming as an explanation enhanced reactivity due to an increased Al2O3 presence (from heating) should review my cited half cell reactions for the Bleach battery detailed previously on the first page of this thread.

[Edited on 30-11-2018 by AJKOER]

AJKOER - 1-12-2018 at 05:05

Pictures of alcohol assisted evaporation displaying mainly amorphous salt(s). Shiny crystals from sunlight reflection likely potassium impurity from sea salt.

Photo Editor-20181201_073507.jpg - 306kB

Photo Editor-20181201_073611.jpg - 352kB

[Edited on 1-12-2018 by AJKOER]

AJKOER - 1-12-2018 at 16:30

Quote: Originally posted by AJKOER

Quote: Originally posted by Texium (zts16)

That's all well and good, but you never proved that YOU made gamma aluminum oxide to begin with!

True, I do not have the means to directly inspect the treated Aluminum and state there exists some gamma aluminum oxide.
......
[Edited on 30-11-2018 by AJKOER]

Actually, there may be some direct evidence that gamma aluminum oxide has been created!

In pictures I have presented relating to the electrosynthesis based attack of heat treated aluminum on SM, I have observed and described in words the creation a white fluffy powder. For example, from http://www.sciencemadness.org/talk/viewthread.php?tid=94166#... :

Quote: Originally posted by AJKOER

Now, at 48 hours, a fluffy white precipitate visible. The aluminum foil displays only minor degradation.
.........
[Edited on 26-9-2018 by AJKOER]

Interestingly, I just came across this reference (link: https://www.researchgate.net/post/What_is_Aluminium_oxide_so... ) to quote Wolfgang H. Muss at Paracelsus Medical University Salzburg:

"gamma-Al2O3 is a hygroscopic powder (white and fluffy) insoluble in water, but in strong acids and bases."

So with the dissolution of the underlying Aluminum metal, any created gamma-Al2O3 may be perhaps visible as a 'white and fluffy' precipitate.

If accurate, this means the technique of heat treating Aluminum metal of equal surface areas/volumes, targeting gamma-Al2O3 formation, may have a visible barometer of success.

--------------------------------------------------------------

More research suggests many possible alumina formation per the experiment from which I referenced the picture, so heat induced gamma alumina formation may not be discernible (see p.39 at https://www.osti.gov/servlets/purl/4037173 ), to quote from this extensive 1960 work:

"Bergmann studied the electrolytic oxidation of aluminum in dilute aqueous hydrogen peroxide and carbon dioxide solutions in order to obtain high-purity aluminas. Using high-purity aluminum as cathode and platinum as anode, he observed the formation of predominantly amorphous products of the composition Al2O3. 2.94H2O to Al2O3. 3.O2H2O containing varying amounts of beta alumina trihydrate (hydrogen peroxide series) and alpha alumina monohydrate-beta alumina trihydrate mixtures (carbon dioxide series), respectively."

What I also find interesting is how Bergman selected an experimental design to create alumina using dilute hydrogen peroxide or aqueous carbon dioxide with a highly noble electrode of platinum (versus my choice of carbon). My selections were derived from my understanding of radical chemistry, which was not significantly advanced at the time Bergman did his experiments.

My understanding is that alpha alumina is similar in appearance to gamma, and that gamma alumina in water in time converts into alpha.

An interesting point on ammonia created clear Al(OH)3 gels (noted on page 69) was that the immediate form of the clear gel is more reactive (like easily reacts with NaOH or HCl), at least for a few hours!

[Edited on 2-12-2018 by AJKOER]

AJKOER - 4-12-2018 at 05:25

Quote: Originally posted by AJKOER

......I am only suggesting, by itself, a temperature treatment path, which per one posited explanation, ascribes an induced gamma alumina layer on larger pieces of Aluminum......

[Edited on 30-11-2018 by AJKOER]

A curious find on the web to quote from Sciencemadness Wiki at http://www.sciencemadness.org/smwiki/index.php/Aluminium_sul...

“Preparation
Aluminium sulfate can be made by reacting sulfuric acid with aluminium oxide, hydroxide, halide or with hot aluminium metal.

2 Al(OH)3 + 3 H2SO4 → Al2(SO4)3 + 6 H2O
2 Al + 3 H2SO4 → Al2(SO4)3 + 3 H2 “

So a possible implication that somehow a hot piece of Aluminum can be attacked by H2SO4 perhaps due to disruption of its protective layer?

Tsjerk - 4-12-2018 at 07:36

https://en.wikipedia.org/wiki/Grimms%27_Fairy_Tales

AJKOER - 4-12-2018 at 11:43

Quote: Originally posted by Tsjerk

https://en.wikipedia.org/wiki/Grimms%27_Fairy_Tales

Perhaps but still don't known, as more research suggests (to me) perhaps the role of solvated electrons in the case of Aluminum alloys, resulting in the hydrogen atom radical (from H+, and some water), which could lead to hydrogen gas disruption of the protective layer (paralleling the gamma alumina and tap water scheme for pure Al), and/or perhaps even a direct sulfate radical attack!

My evidence, the property of Aluminum ALLOYS to apparently rapidly attack Al if H2SO4 concentration is between 40% to 95% with a peak at 80%, but not so for pure Al, as that requires heating the aluminum metal.

In the case of Iron and H2SO4, the process is one of 'dissolution and diffusion' of any formed sulfate layer away from the surface.

Source: See https://books.google.com/books?id=KXwgAZJBWb0C&pg=RA1-PT... .

[Edited on 4-12-2018 by AJKOER]

j_sum1 - 4-12-2018 at 14:56

Quote: Originally posted by AJKOER

My evidence, the property of Aluminum ALLOYS to apparently rapidly attack Al if H2SO4 concentration is between 40% to 95% with a peak at 80%, but not so for pure Al, as that requires heating the aluminum metal.

I would love to see a write-up of the experiment you did on this. My (limited) experience is that Al is pretty much impervious to acids unles there are halides around or something like Hg or Ga to weaken the structure at the grain boundaries.
Specifically, what kind of alloy are we talking about here that is attacked by H2SO4? And how much different is it from the behaviour of pure Al?

It would be nice to establish some basic observations before speculating on complex mechanisms.

AJKOER - 4-12-2018 at 16:02

Quote: Originally posted by j_sum1

Quote: Originally posted by AJKOER

I usual quote from my source, but not doable in the 2 pages presented commentary in the 'Handbook of Corrosion Data' which refers to what it calls 'Material Summaries' from 'a survey of the available literature'. In the 'Aluminum.' section its basically generally states what I say with respect to Al alloys used industrially. Also, there is some detail with respect to certain Aluminum alloy numbers, etc.

As to what literature, what precise alloys, date of various works,..., someone can spend $618.72 for details, if included, as I suspect this expert book is likely for people making decisions involved in designing or buying containers to hold/transport various acids and such and avoiding significant corrosion issues (hence the value added book price).

This reference focuses on industrial use of Al (and other metals) expressed in the form of available numbered alloys (one of which may be the pure metal).
---------------------------------------

Some background points on the Kinetics of Corrosion:

1> Alloys with a highly electropositive metal like Aluminum may create a galvanic corrosion scenario starting with the flow of electrons.

2> Next, to quote a source, ‘Kinetics of Corrosion Inhibition of Aluminum in Acidic Media by Water-Soluble Natural Polymeric Pectates as Anionic Polyelectrolyte Inhibitors’ by Refat M. Hassan and Ishaq A. Zaafarany, in Materials (Basel). 2013 Jun; 6(6): 2436–2451, doi: 10.3390/ma6062436, link: https://www.ncbi.nlm.nih.gov/pmc/articles/PMC5458939/ :

“The corrosion rate was found to be a function of the concentration of the acid. This result indicates that at least one of the corrosion paths of dissolution of Al metal in HCl solution should involve the presence of hydrogen ions in the rate-determining step…
3.1. Corrosion Mechanism
We propose a suitable mechanism of corrosion, in accordance with the above experimental observations. The corrosion of metal involves an electrochemical process [47,48,49] resulting from dissolution of Al metal in the acid. This process can be expressed by the anodic and cathodic processes, which are defined by Equations (4) and (5), respectively,
Al(s) <--Ox--> Al3+ + 3 e- (4)
2 H+ + 2 e- <--Red--> H2 (5)
The overall electrochemical process can be written as follows:
(6) 2 Al(s) + 6 H+ <-->2 Al3+ + 3 H2(g) (6)
The cathodic reaction produces Hchemisorbed by picking up an electron that released in the anodic reaction (H+ + e = Hchemisorbed ) in Al corrosion in HCl. In such acidic solutions, the Hchemisorbed on the metal surface reacts by combining with other adsorbed Hchemisorbed to form H2 gas molecule, which bubbles from the metal surface. A very small amount of the uncombined Hchemisorbed will remain; however, this amount does not affect the whole process. “

3> Also from a source: ‘The mechanism of hydrogen embrittlement: the stress interaction between a crack, a hydrogen cluster, and moving dislocations’ by A.T. Yokobori Jr. , In International Journal of Fracture, July 2004, Volume 128, Issue 1–4, pp 121–131, link: https://link.springer.com/article/10.1023/B:FRAC.0000040974.... , to quote:

“In this paper, the stress corrosion cracking model on the basis of hydrogen diffusion and concentration toward the elastic-plastic stress field around a crack and the interaction of dislocations and hydrogen around a crack tip are proposed to clarify the mechanism of stress corrosion cracking for ductile and brittle materials.”

Taking into account the above, my summary on the removal of the protective layer begins with electrons reacting with H+ to form the hydrogen atoms radical. There is then a the diffusion of hydrogen atoms radicals which have been chemisorbed. With increasing concentration of the hydrogen atom radical, some combining creating hydrogen gas bubbles. This can lead to cracking and dislodging of the aluminum oxide layer.

[Edited on 5-12-2018 by AJKOER]

j_sum1 - 4-12-2018 at 16:49

I ask because you said, "my evidence". I assumed you had done some kind of experiment.

It should not be too difficult to cite some literature if you have a good source. It does not need to be pages and pages. In fact, one good quotation and maybe a cut and paste of a supporting diagram is probably better.

You seem to be claiming a number of unfamiliar properties of aluminium. I confess to not being able to follow all that you have written. Before a deep discussion of mechanisms and radicals I think establishment of some fundamental properties is important. I am not aware of any significant reaction between alloys of Al and sulfuric acid. I may well be wrong. I am asking you to fill in that blank before taking us deep down the rabbit hole.

[edit]
You have added some material that was not present when I composed my post.

Among your references I see some standard electrolysis oxidation and reduction.
I also see something about stress corrosion cracking.
Neither of these seem pertinent to the discussion (which originally was acetates and cyanides). I am now challenging you to back up what you have claimed with respect to sulfuric acid.

It seems to me pointless to make such a large number of tangential claims and not connect them to an overall objective.
Likewise it seems pointless to write in such a convoluted fashion that others on the board (smarter people than me) cannot follow.
And it seems strange that you do not respond to specific challenges to the data you present and the claims you make.

[Edited on 5-12-2018 by j_sum1]

AJKOER - 4-12-2018 at 17:04

A more readable reference with graphs, see https://books.google.com/books?id=iEeiQEeLOmYC&pg=PA33&a... which does confirm generally that pure Al is more corrosion resistant than an Aluminum alloy. Interestingly, in Figure 7 the author notes that aluminum oxide containing iron is actually a semiconductor, capable of allowing the passage of electrons, and likely fostering the formation of local galvanic cells (leading to corrosion).

Also discussion at https://books.google.com/books?id=iEeiQEeLOmYC&pg=PA39&a... . In particular, read paragraphs relating to "Effect of Purity: Trace Elements" and "Effect of Hydrogen" on that page.

Of note, once corrosion is present, apparently in some cases, the rate of corrosion can proceed proportional to the concentration of NaCl (an accelerant).
-------------------------------------

With respect to H2SO4 specifically, see Figure 3, p. 689, at https://www.nrc.gov/docs/ML0633/ML063390144.pdf that clearly shows at around 80%, a contour displaying the highest corrosion in inches per year of 1.2 . To quote a general comment from the text:

"Dilute sulfuric acid solutions, up to ~ 10% in concentration, causes some attack on aluminum base alloys, but the action is not sufficiently rapid at room temperature to prevent their use in special applications. In the concentration range of ~40 to 95%, rather rapid attack occurs. In extremely concentrated or fuming acid, the rate of attack drops again to a very low value."

Also:

"As evident in Table 8, aluminum (and its alloys) becomes the anode in galvanic cells with most metals, protecting them by corroding sacrificially. Only magnesium and zinc are more anodic and corrode to protect aluminum. This type of corrosion can be found in strong acidic or strong basic solutions, as illustrated in Figure 6 [4,1O]. The rate of corrosion can vary from several microns per year to several microns per hour."

Also with respect to a temporarily more reactive Al(OH)3 gel created by ammonia precipitation subject to time aging, the comment:

"The aluminum hydroxide gel is not stable, but crystallizes with time to give, first, the rhombohedral monohydrate (Al2O3-H2O or boehmite), then the monoclinic trihydrate (A12O3.3H2O or bayerite), and finally another monoclinic trihydrate (hydrargilite). This development of aluminum hydroxide is known as "aging" [4]."

Here is a guide with another source as to explaining what could be occurring with regard to varying acid concentrations acting on Aluminum, to quote (see http://www.conways.co.za/pdf/afsa_corrosion_pocket_guide.pdf ):

“In the case of aluminium and other non ferrous metals, this sulphuric acid is absorbed in the corrosion process. It gives rise to the formation of metallic sulphates. However, in iron and steel the sulphuric acid is regenerated as a result of hydrolysis and the corrosion reaction continues.”

So lower than 100% H2SO4 concentrations could be engaging with water in a recycling action on the attack of the protective layer for select Aluminum aollys (with Fe,...). A peak concentration of 80% could be reflective of the significance of the recycling action in the presence of water. I would think that water presence for creating solvated electrons could also be a part of the process.

Also, with respect to precise behavior of H2SO4 at varying concentrations, note this commentary in the case of its action on steel (not aluminum):

“Depending on acid concentration and the metal involved, sulfuric acid can behave as either a reducing acid or an oxidizing acid. In lower concentrations (up to about 70%) and temperatures on steel or stainless steel, it is reducing. At concentrations over 70%, and into the oleum range, the acid is oxidizing to steel and stainless steel.”

where, with respect to Aluminum metal, I would expect a shift to the oxidizing concentrations of H2SO4 would be less corrosive, so 80% appears reasonable with high H+, 20% water for possible recycling path to H+ and solvating electrons, and not deep into the oxidizing range commencing at 70%.

[Edited on 5-12-2018 by AJKOER]