AJKOER
Radically Dubious
Posts: 3026
Registered: 7-5-2011
Member Is Offline
Mood: No Mood
|
|
Dawsonite via the Aluminum air battery using Na2CO3 electrolyte?
First we note on heating up aqueous Na2CO3 (38 grams heated in 50 ml of 3% H2O2 in microwave), we have the reaction:
2 Na2CO3 + 2 H2O = 2 NaHCO3 + 2 NaOH
which on boiling, in the presence of Aluminum, is known to be rapidly attacked.
Also, possibly to a more limited extent as well the breakdown of the aqueous Sodium bicarbonate on boiling (although the supplied reference claims
completion on boiling):
2 NaHCO3 + H2O ---Boiling → CO2(g) + 2 H2O + Na2CO3 (see http://antoine.frostburg.edu/chem/senese/101/inorganic/faq/c... )
Or, on net, an optimistic limiting and most likely inaccurrate overall reaction even in the presence of a reactive metal like Aluminum to move the
reaction to the right:
Na2CO3 + H2O ---Boiling → CO2(g) + 2 NaOH
And indeed a limiting one as the following reaction has come to my attention, in the presence of say Al(OH)3, consuming the NaHCO3. To quote https://books.google.com/books?id=FijWBgAAQBAJ&pg=PA149&...
Al(OH)3 + NaHCO3 = NaAl(CO3)(OH)2 + H2O
Now, we bring into play some electrochemistry, to quote Wikipedia on the Aluminum-air battery (https://en.m.wikipedia.org/wiki/Aluminium–air_battery#Elec... :
"The anode oxidation half-reaction is Al + 3OH− → Al(OH)3 + 3e− +2.31 V.
The cathode reduction half-reaction is O2 + 2H2O + 4e− → 4OH− +0.40 V.
The total reaction is 4Al + 3O2 + 6H2O → 4Al(OH)3 + 2.71 V.
About 1.2 volts potential difference is created by these reactions, and is achievable in practice when potassium hydroxide is used as the electrolyte.
Saltwater electrolyte achieves approximately 0.7 volts per cell."
Note the overall consumption of water which could have important implications for pH and solubility.
Now, in our case upon adding Aluminum, to boiled Na2CO3 in the presence of oxygen (or an oxygen source like H2O2 employed here), we have a better
electrolyte than saltwater but not as effective as the KOH. Note, it is O2 and H2O attacking the Al and not the electrolyte!
At cooler temperatures (under 50 C) and agreeable pH, some Sodium aluminate could form via the equilibrium reaction:
Al(OH)3(s) + NaOH = NaAl(OH)4
However, as the Na2CO3 electrolyte was not heated to boiling for more than 30 seconds in a microwave at start of this experiment, it for the most part
should still be present. And given the reactions:
Na2CO3 + H2O = NaHCO3 + NaOH
Al(OH)3 + NaHCO3 = NaAl(CO3)(OH)2 + H2O
I suspect that the major product formed is, in fact, Dawsonite (see https://en.m.wikipedia.org/wiki/Dawsonite ) confirmed by coloring and texture distinct from my usual Al(OH)3 suspension. I also extracted some
clear electrolyte which could be aqueous NaAlO2 and placed drops into a dilute vinegar solution. What I observed was not white clouds, but small thick
globs (like those sitting in the reaction vessel) that now formed bubbles with acetic acid, indicating a possible carbonate salt (see last photo).
To speed things up in the electrochemical half reactions, I used 50 ml of 3% H2O2 as an oxygen source and just one microwave heating for 30 seconds.
My Aluminum source was standard home Al foil in the size of 13.5 cm X 12 cm, which was further briefly heated in a methane flame to remove protective
coating and induce the formation of a weakened gamma Al2O3 on the foil, which enhances the reactivity, per my many observations, as was suggested in a
source. Here is a link to a white paper on the topic https://www.google.com/url?sa=t&source=web&rct=j&...
Pictures below show the entire mass of Aluminum foil at start, middle and end being nearly completely consumed in under 1O minutes, no added heating
applied.
  
[Edited on 29-7-2015 by AJKOER]
|
|
|
Oscilllator
National Hazard
Posts: 659
Registered: 8-10-2012
Location: The aqueous layer
Member Is Offline
Mood: No Mood
|
|
Although your product might have the same formula as dawsonite, I doubt it will be very easy to get it to form the same crystal structure, for the
same reason that burning iron in sulfur will not get you nice crystals of pyrite.
|
|
|
AJKOER
Radically Dubious
Posts: 3026
Registered: 7-5-2011
Member Is Offline
Mood: No Mood
|
|
Quote: Originally posted by Oscilllator  | | Although your product might have the same formula as dawsonite, I doubt it will be very easy to get it to form the same crystal structure, for the
same reason that burning iron in sulfur will not get you nice crystals of pyrite. |
A fair point, I am no expert in the manufacture of rare minerals.
Interestingly, per a US Patent 4,221,771, I recently found (link: http://www.google.com/patents/US4221771 ), to quote:
"Dawsonite is formed by reacting alumina trihydrate with sodium bicarbonate in a sodium carbonate solution under pressure at a temperature between
100° and 200° C. for 30 minutes to 4 hours."
My cited preparation mirrors this with the differences that my Al(OH)3 is prepared in situ in a microwave assisted process (known to reduce processing
time and create more uniform and smaller nano-suspensions, which are inherently more reactive, see references below), the source of the NaHCO3 is from
the hydrolysis of the Na2CO3 input, and finally, no pressure or prolonged heating appears to be necessary in my preparation.
One of the evident issues, however, with my procedure is that the purity (or problematic impurity) of the starting Aluminum metal source used to
produce the Al(OH)3 may presence a quality issue.
---------------------------------------
A reference supporting my comment: "Microwave-Assisted Chemistry: Synthetic Applications for Rapid Assembly of Nanomaterials and Organics", to quote:
"Microwave-assisted chemical reactions are now well-established practices in the laboratory setting although some controversy lingers as to how MW
irradiation is able to enhance or influence the outcome of chemical reactions. Much of the discussion has focused on whether the observed effects can,
in all instances, be rationalized by purely thermal Arrhenius-based phenomena (thermal microwave effects), that is, the importance of the rapid
heating and high bulk reaction temperatures that are achievable using MW dielectric heating in sealed reaction vessels, or whether these observations
can be explained by so-called “nonthermal” or “specific microwave” effects."
Link: http://pubs.acs.org/doi/abs/10.1021/ar400309b
Here is a related example of how MW radiation may be augmenting the formation of ferrates via the presence of Fe(OH)3. See, for example,
"Nanostructures: Synthesis, Functional Properties and Application" edited by Thomas Tsakalakos, ..., pages 56 and 57 at https://books.google.com/books?id=z2ryCAAAQBAJ&pg=PA57&a...
To paraphrase the authors, apparently, Ferric hydroxide is very good at absorbing electromagnetic radiation and converting it into a homogeneous and
effective heating source. In a nano-suspension of Fe(OH)3, this can increase nucleation sites and limit average particle growth while narrowing the
particle size distribution. This is attributed to increase kinetics and more homogeneous heating resulting in an increase in yield and reduced
processing time. In general, magnetic materials under MW heating appear to be superior as a consequence of a ferromagnetic resonance effect.
[Edited on 29-7-2015 by AJKOER]
|
|
|
blogfast25
International Hazard
Posts: 10562
Registered: 3-2-2008
Location: Neverland
Member Is Offline
Mood: No Mood
|
|
All the references in the world do not change the FACT that you have no real proof you've prepared Dawsonite, AJ.
Try preparing more of that material, wash exhaustively to get rid of soluble carbonate and determine composition of the dried product. Or at a very
minimum CO<sub>2</sub> content.
Wiki's (and other sites) images of Dawsonite suggest a material that was produced geologically via hydrothermal route. It's unlikely you've created
the conditions to prepare NaAlCO<sub>3</sub>(OH)<sub>2</sub>. If preparing such a compound is possible on simple lab scale,
we'd have heard about it by now. Patents are poor and often biased sources of information.
[Edited on 29-7-2015 by blogfast25]
|
|
|
AJKOER
Radically Dubious
Posts: 3026
Registered: 7-5-2011
Member Is Offline
Mood: No Mood
|
|
Blogfast:
Your point is in order as the author of the cited patent has clearly tested other existing patents claiming the mineral, and noted, in his assessment,
at least one instance of where, per his chemical depiction of Dawsonite as represented by the equation below, to quote:
" 2 Al(OH)3 +2 NaHCO3 =Na2O:Al2O3 :2 CO2 :2 H2O + 2 H2 O "
there is at least one other witnessed possible chemical composition. For example, to quote, "compounds described in U.S. Pat. No. 3,115,387 (Al2O3
:H2O:NaHCO3) have a composition different from dawsonite.") However, as my outlined path is, if anything, deficient in NaHCO3 based on the hydrolysis
of Na2CO3, I feel this is not too likely.
Also, given the seeming unintended parallel between my path and US Patent 4221771 as I detailed above, I am feeling good about the prospects, and if
necessary, modifications in my ratio of starting reagents, and/or conditions, to achieve the above agreed upon composition.
I also feel more comfortable with the quality of this patent as is perhaps to be expected given prior patents on Dawsonite itself and some apparent
commercial significance to quote:
"Dawsonite and the other alkali metal alumino carbonates are useful in several applications including, for example: in antiacid compositions, as an
antiperspirant, as a fire extinguishing powder, as an ingredient for self-extinguishing plastics, as an ingredient for fire retardant intumescent
paints and as a blowing agent."
---------------------------------------
Note, twice my specified formula in the opening thread for Dawsonite, namely, 2NaAl(CO3)(OH)2 equals Na2O.Al2O3.2CO2.2H2O as cited per the patent
above.
[Edited on 29-7-2015 by AJKOER]
|
|
|
blogfast25
International Hazard
Posts: 10562
Registered: 3-2-2008
Location: Neverland
Member Is Offline
Mood: No Mood
|
|
The patent you first cited requires about 9 atm pressure.
Prepare more of that particular precipitate and at a minimum determine CO<sub>2</sub> content on the clean, dry product. That's not hard
to do.
My money is on not finding any, BTW... But it would be interesting if your precipitate did contain some bound carbonate.
[Edited on 29-7-2015 by blogfast25]
|
|
|
AJKOER
Radically Dubious
Posts: 3026
Registered: 7-5-2011
Member Is Offline
Mood: No Mood
|
|
My sample has changed today appearing as a fine white powder that after exhaustive washing, no longer fizzles with vinegar.
As such, I plan to perform some more stoichiometric calculations based on my equations and review other cited patents for guidance on proper
proportions and conditions.
Interestingly, while not likely, there may also be some commercial significance to my suggested preparation, so I will understand if the room goes
silent.
[Edited on 29-7-2015 by AJKOER]
|
|
|
aga
Forum Drunkard
Posts: 7030
Registered: 25-3-2014
Member Is Offline
|
|
It tends to go silent when $ are involved, as most people don't ever get that far - turning an idea into $.
Nice idea though AJ. Like the microwaves.
|
|
|
Oscilllator
National Hazard
Posts: 659
Registered: 8-10-2012
Location: The aqueous layer
Member Is Offline
Mood: No Mood
|
|
| Quote: Originally posted by AJKOER | | My sample has changed today appearing as a fine white powder that after exhaustive washing, no longer fizzles with vinegar. |
Going from a clear, crystalline substance to a white powdery one is often an indicator that some water of hydration has been lost. If you have a nice
sensitive scale, you might want to test if that's the case.
|
|
|
|
![[*]](images/xpblue/default_icon.gif){kind=icon}
