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kilowatt
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Ammine Aluminum Complex
Does aluminum readily form any ammine complexes in aqueous solution? I found just a few google hits for hexamminealuminum, but only under rather
harsh conditions. I'm considering making aluminum nitrate and ammonia from electrolysis in AN solution (or perhaps just aluminum oxides/hydroxides in
AN) and don't want to run into any problems.
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woelen
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As far as I know, aluminium does not form ammine complexes. If you add aqueous ammonia to aluminium salts in solution, then a precipitate of Al(OH)3
is formed, which does not redissolve on addition of more ammonia. This separates aluminium from metals like copper and zinc.
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12AX7
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I would suppose you might get something dissolving AlCl3 in liquid NH3. The reaction should be analogous to H2O, but milder due to the weaker nature
of NH3 and lower temperature (unless ran under pressure).
In aqueous solution, I don't see it happening.. aluminum *really loves* oxygen, even if it has hydrogens stuck on as in water.
Tim
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kilowatt
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I am not trying to make an amminealuminum nitrate complex; rather I am trying *not* to, looks like that's covered though. Now as for the aluminum
nitrate formation, it should go forward as long as ammonia is removed from the system correct? I imagine aluminum hydroxide would precipitate and
then be dissolved as the nitrate content is enriched past that of ammonium by electrolysis.
[Edited on 6-8-2008 by kilowatt]
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woelen
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You won't get aluminium nitrate (assuming you are electrolysing with an Al-anode). At the cathode, you get H2 and NH3. At the anode, Al goes in
solution as Al(3+). With the NH3, and water from the solution, this forms Al(OH)3 and ammonium ion.
Only if you find a way to remove ammonia you'll get aluminium nitrate, but the ammonia does not escape as gas, only the H2 escapes as gas.
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kilowatt
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What about a membrane that is much more permeable to ammonium ions than aluminum? Is there such a thing? I am really after one that is effective at
blocking lead ions.
Ammonia has rather low solubility in water up around 80-100°C anyway. If ammonia wasn't being removed from these ammonium nitrate electrolysis
systems, you would not be able to smell it very strongly or see it bubbling out. However, you can in practice, especially when the cell has heated
up.
[Edited on 7-8-2008 by kilowatt]
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Rosco Bodine
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Depending on the solubility of NH4NO3 in DMSO, as compared to the solubility of anhydrous Al(NO3)3 in DMSO, you may be able to get somewhere using
amalgamated aluminum, with evolution of both ammonia and hydrogen, and conversion of the Al to anhydrous Al(NO3)3 .
This might also occur in one of the eutectic melts of
NH4NO3 with other materials, including the possibility
that a tertiary eutectic involving or derived from NH4NO3 and Al(NO3)3 formed by this route may be the reaction result.
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kilowatt
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What's the point of using amalgam instead of solid aluminum?
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Rosco Bodine
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Because the amalgam breaks the passivating oxide layer on the Al surface and allows the elemental aluminum dissolved in the advancing film of mercury
to react , continually replenishing fresh elemental aluminum to the reaction front, using the film of mercury as a conveyor.
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kilowatt
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But the mercury would be nitrated too. Also the nitrate ions should break the oxide layer easily.
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Rosco Bodine
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No the mercury would not be nitrated too, at least except as a possible intermediate and regenerated catalyst. You see, if it was, it would
immediately reverse react itself being reduced again to the metallic mercury by the more active aluminum. And no, the nitrate would not break the
oxide, but would in fact agressively passivate the aluminum completely.
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DJF90
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Rosco is right, aluminium is way above mercury in the reactivity series  However
the use of mercury is not something I would like to deal with as clean-up/disposal is a real PITA. But it is definately a viable option.
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kilowatt
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Oh. It would still have to be done in DMSO or a similar solvent though right? I lack that and it's synthesis would be tedious (at least the way I
imagine it would be done). Any less exotic solvents that might work? What properties are we after here? Just something that doesn't react with
aluminum?
| Quote: |
And no, the nitrate would not break the oxide, but would in fact agressively passivate the aluminum completely. |
Think it would work for hardcoat anodizing?
Oh by the way, I was looking at aluminum as a less toxic substitute for my lead nitrate process for making nitric acid. The use of mercury sort of
shoots that one to hell.
[Edited on 7-8-2008 by kilowatt]
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Rosco Bodine
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Mercury is useful and most of it is recoverable. It stays attached to the unconsumed aluminum. DMSO is a veterinary item available at feed stores and
used as a
liniment on horses and sometimes on the horses owners
The idea I had there with DMSO is a non-aqueous solvent which could (possibly) dissolve the NH4NO3 and would (hopefully) not react with the aluminum
either or react with the anhydrous aluminum nitrate desired to be formed. Anhydrous conditions would be needed for avoiding the formation of basic
aluminum nitrate and/or hydrated forms of aluminum nitrate. Alcohol couldn't be used because it would form an aluminum alkoxide.
Awhile back, there was a look taken (by me) at some schemes for use of dehydrating nitrate salts as a possible way of enhancing nitration mixtures,
and some of those double salts might be more useful and easier to make than Al(NO3)3 .
http://www.sciencemadness.org/talk/viewthread.php?tid=4701&a...
[Edited on 8-8-2008 by Rosco Bodine]
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kilowatt
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Thanks for the info Rosco. I might just try the amalgam. I like to be able to recover *all* the mercury though; it is rather hard to obtain and I
only have a couple kilos most of which is reserved for my amalgam cell. Thanks for the old thread too.
Would DMSO react with nitric acid formed?
What if, instead of using DMSO, I used a tyvek membrane to separate an anode chamber containing water from a cathode chamber containing ethanol? I
know at least lead nitrate is only minimally soluble in ethanol, but an electrolyte such as calcium chloride could be dissolved in it. Other solvents
may work even better.
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Rosco Bodine
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I would highly recommend alternately using magnesium salts since there is some information already available to use as a guide. However, the matter
of aluminum amalgam is intriguing, especially in an electrolyte or reaction system containing NH4NO3, because of the
possible formation of ammonium amalgam. It is a huge unknown what may result, and a low voltage AC across
two aluminum amalgam electrodes might be an interesting experiment also.
I don't follow your reasoning on the divided cell, and I am not sure that nitric acid will even be formed at all, but with regards to the DMSO, that
is an unknown also. Amalgams of aluminum and other passivating metals can be extremely reactive, very powerful reducing agents, and the
tendency of mercury to escape would be minimal if you think about it. That is precisely why people can have
dental amalgams in their bodies without being poisoned by the mercury, it remains sequestered and tied to the
amalgam alloy in preference to just floating off into space
as ions or salts.
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kilowatt
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I'd have to use the magnesium as oxide or amalgam to make it practical since magnesium is relatively difficult to isolate in its pure form. As a
relatively expensive metal it would definitely have to be a cyclic process. Resultant magnesium oxide from the decomposition step would have to be
chlorinated or otherwise prepared for dissolution and electrolyzed into amalgam to recover. Will I be able to use water as an electrolyte?
Magnesium nitrate has a higher decomposition temperature than aluminum nitrate making it less desirable, however it will work especially if the rest
of the chemistry is easier.
My idea with the divided cell is to prevent the metal cations from entering the catholyte and being removed from solution at the cathode.
[Edited on 9-8-2008 by kilowatt]
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kilowatt
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I think I will use the magnesium oxide + AN method; it is more suitable than any other method I have conceived, and cyclic. I am doing a test run
right now; it seems to be working very well. Very high molarity per weight on the magnesium salt too. Thank you Rosco!!
[Edited on 16-8-2008 by kilowatt]
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Rosco Bodine
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You could just use epsom salt as a cheap source for magnesium, and precipitate as the carbonate using baking soda, and then react the carbonate with
NH4NO3.
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DerAlte
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Isn't magnesium hydrogen carbonate soluble, like the Ca salt? Not very, perhaps. Better to use sodium carbonate than Baking soda, I think - or boil
the bicarbonate first, or heat it.
Der Alte
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Rosco Bodine
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It's been years since I did this precipitation.
Boiling hot solutions are part of the process I was contemplating. The washing soda does work, the product being "magnesia alba".
Magnesia alba, a white bulky precipitate obtained by adding sodium carbonate to Epsom salts,is a mixture of Mg(CO 3 H) (OH) 2H 2 O,Mg(CO 3 H) (OH) and
Mg(OH) 2. It is almost insoluble in water, but readily dissolves in ammonium salts.
http://books.google.com/books?hl=en&id=i6A6AAAAMAAJ&...
http://books.google.com/books?id=ow8AAAAAQAAJ&pg=PA222&a...
Gmelin states that the tenacity with which magnesium nitrate
holds water is evident in that the monohydrate remains intact even at the melting point of lead !!!! It is also remarkable that the nitrate salt
itself would even remain intact at so high a temperature.
[Edited on 17-8-2008 by Rosco Bodine]
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not_important
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| Quote: |
Gmelin states that the tenacity with which magnesium nitrate
holds water is evident in that the monohydrate remains intact even at the melting point of lead !!!! It is also remarkable that the nitrate salt
itself would even remain intact at so high a temperature. |
Silver nitrate is stable up to near 440 C, well above lead's melting point. A standard purification method is to melt AgNO3 and keep it molten for a
few minutes, cool, dissolve in DW. and filter; copper, iron, and similar nitrates will have decomposed to oxides.
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kilowatt
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I wonder if calcium oxide, perhaps better yet strontium oxide, would react faster than that of magnesium. Precipitating a carbonate would use up too
many other resources unless I synthesized it using CO2 which is sufficiently cheap, but more difficult. Using calcium oxide itself, it tends to react
with water forming the hydroxide.
| Quote: |
Silver nitrate is stable up to near 440 C, well above lead's melting point. A standard purification method is to melt AgNO3 and keep it molten for a
few minutes, cool, dissolve in DW. and filter; copper, iron, and similar nitrates will have decomposed to oxides. |
For purifying Ag, why not just add a little NaCl to the AgNO3? That method instantly precipitates AgCl which can be heated to obtain 0.999 pure
silver. The other metals of course remain soluble.
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Rosco Bodine
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You seem to be wanting to make things difficult.
Washing soda, sodium carbonate is dirt cheap.
Arm and Hammer washing soda in the yellow box
is a hydrate, lower than the decahydrate...I did a weight loss on baking one time to determine what it was and reported the crude analysis here in
another thread.
The off the shelf washing soda which I baked out,
turned out to be 83.4% Na2CO3 , and the rest water of crystallization.
http://www.sciencemadness.org/talk/viewthread.php?goto=lastp...
[Edited on 17-8-2008 by Rosco Bodine]
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kilowatt
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Nah if wanted to make things difficult there's a ton of complicated methods which I could use, most of which have been discussed and thrown out.
Magnesium salts would have never even occurred to me if it wasn't for your input, since my train of thought was on aluminum or transition metal salts.
There just doesn't seem to be a method suitable by using them; alkaline earth salts seem to offer the best solution.
Magnesium oxide reacts suitably (though slowly) with NH4NO3, so I don't see the point in using additional chemicals to treat it first. I'm looking to
process over 200 moles (perhaps much more) of nitrate here one way or another so it really adds up and really is prohibitive if you are using excess
stuff that is not conveniently cyclic. I know of no simple way to prepare magnesium carbonate from magnesium oxide (the magnesium source I am using
and will be re-using after the decomposition of the nitrate) anyway. To prepare more soluble magnesium salts to do the carbonate precipitation I
would have to react the oxide with mineral acid. I do have plenty of sodium bicarbonate which I could easily decompose to carbonate, but would the
extra steps and use of material be enough to make it worth it? I also have CO2 and could prepare the carbonate directly from the oxide/hydroxide with
that, but that is a slow process. How much would the final reaction rate even be increased? Surely not enough to make up for the additional steps.
Of course I do have some magnesium chloride hexahydrate (quite cheap, sold as ice melt) that I have been considering converting, and for the first
reaction pass I could precipitate it with sodium carbonate and definitely come out ahead. However I don't have enough to process the amount of
nitrate I am wanting to, nor do I want to expend that much sodium carbonate (could make more with the Solvay Process of course, since what I am doing
produces plenty of ammonia which I would ultimately like to store in solid form such as NH4Cl). I will be post back soon with a comparison of
magnesium oxide and carbonate for the reaction since the next thing I'm going to do is precipitate some MgCO2.
| Quote: |
Isn't magnesium hydrogen carbonate soluble, like the Ca salt? Not very, perhaps. |
It is considerably soluble, but much less soluble than sodium carbonate, magnesium sulfate (epsom salt) or magnesium chloride. Thus the precipitation
will work with a minimal amount of water.
[Edited on 17-8-2008 by kilowatt]
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