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.woelen - 5-8-2008 at 22:23
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.12AX7 - 6-8-2008 at 07:05
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.
Timkilowatt - 6-8-2008 at 09:41
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]woelen - 6-8-2008 at 22:30
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.kilowatt - 7-8-2008 at 09:56
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]Rosco Bodine - 7-8-2008 at 12:12
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.kilowatt - 7-8-2008 at 15:22
What's the point of using amalgam instead of solid aluminum?Rosco Bodine - 7-8-2008 at 15:44
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.kilowatt - 7-8-2008 at 15:54
But the mercury would be nitrated too. Also the nitrate ions should break the oxide layer easily.Rosco Bodine - 7-8-2008 at 16:06
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.DJF90 - 7-8-2008 at 16:32
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.kilowatt - 7-8-2008 at 17:16
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]Rosco Bodine - 7-8-2008 at 19:27
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]kilowatt - 8-8-2008 at 17:57
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.Rosco Bodine - 8-8-2008 at 19:31
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.kilowatt - 8-8-2008 at 23:32
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]kilowatt - 16-8-2008 at 19:47
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]Rosco Bodine - 16-8-2008 at 20:52
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.DerAlte - 16-8-2008 at 21:43
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 AlteRosco Bodine - 16-8-2008 at 23:26
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.
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]not_important - 17-8-2008 at 06:33
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.kilowatt - 17-8-2008 at 12:32
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.Rosco Bodine - 17-8-2008 at 14:40
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.
[Edited on 17-8-2008 by Rosco Bodine]kilowatt - 17-8-2008 at 15:50
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]Rosco Bodine - 17-8-2008 at 16:30
What is the purpose of the process? It may help me to understand what you are trying to accomplish.kilowatt - 17-8-2008 at 17:14
To produce nitric acid and ammonia from ammonium nitrate without using up any other reagents or materials. For example react magnesium oxide to
ammonium nitrate to produce ammonia and magnesium nitrate, then decompose the nitrate to yield nitric acid synthesis gases and magnesium oxide, which
is used again in the next reaction pass. For the first reaction pass I could use faster methods like the magnesium carbonate, which I have enough
materials to cheaply produce over 100 moles of. However I am interested in making far more nitric acid than that, so I will need a more cyclic route
(like using the MgO directly) to continue.
I was going to use aluminum with electrolytic methods, hence the title of this thread (I had an older thread about using copper), however magnesium is
obviously a better choice.
The nitric acid will be in turn used to make lead nitrate from scrap, thus separating the alloy since tin will form oxide and precipitate (its nitrate
decomposes in water). Antimony should precipitate as well. Lead nitrate can then be decomposed to regain the nitric acid, and leave lead oxide for
smelting.
Magnesium nitrate itself should also be an excellent oxidizer for solid rocket propellants.Rosco Bodine - 17-8-2008 at 17:54
Decomposing Magnesium Nitrate will require very high temperatures, and recycling the decomposition gases back to nitric acid may not be so
straightforward as
reacting with water.
If experiments with producing nitric acid more efficiently
using the nitrate/sulfuric acid reaction is a goal, getting
2 moles of nitric acid per mole of sulfuric, avoiding the
bisulfate byproduct ....then magnesium nitrate makes sense.
If experiments with alternate dehydrating agents for nitrolysis and nitration is a goal, then magnesium nitrate
or the double salt with ammonium nitrate, in solution in strong nitric acid makes sense.
But if making lead nitrate is the goal, then there are probably easier ways, already outlined in the preparation of lead salts thread, and for a
scheme involving ammonium nitrate as the nitrate donor magnesium nitrate
as an added step seems more trouble and no advantage,
which doesn't make sense.
Did you try low voltage ac on lead scrap or amalgamated
lead scrap in moderately hot aqueous NH4NO3?
The frequency of the AC may have to be lowered to
something like one cycle every two minutes, using a
double pole double throw relay driven by a timer to
make square wave AC from DC. What I was thinking
is to cycle it at a rate which allows formation and saturation of an amalgam film with ammonium amalgam, decomposed on reversing polarity in the warm
solution.
It might work even without an amalgam, or it may require a divided cell or a pumped liquid mercury electrode with
the ammonium amalgam decomposed in a separate compartment. But these things are definitely worth looking at if you are processing large amounts of
scrap lead.
Some of the straight chemical methods using HCl as a cheap acid source and a nitrate would probably also be more economical than any route using
magnesium nitrate.
I just don't see the recycling loop you are contemplating there as being viable.
Using HNO3 gotten from a distillation of HCl plus Magnesium Nitrate is another possibility which might be worth looking at....I'm not sure how
that particular reaction efficiency would work out if at all. But for sure the conversion of even low concentration H2SO4 like battery electrolyte
reacted with Magnesium Nitrate would
provide a pure distilled HNO3, and the byproduct Magnesium Sulfate could be recycled via Magnesia Alba
and Ammonium Nitrate......the only byproducts of the loop
would be glauber salt and ammonia. If this is a battery recycling project, that provides utilization of the battery electrolyte as an active part of
the recycling, eliminating any need for its disposal by using it in the recycling of the lead and lead oxide and sulfate.
If thermal decomposition gases from the lead nitrate are neutralized by alkali, you will get a mixture of nitrates
and nitrites. Or you could use those "nitrous gases" for other interesting purposes, much more easily than recycling them to nitric acid. Plan on a
catalytic converter
being part of the process if recycling the nitric acid is desired.
[Edited on 17-8-2008 by Rosco Bodine]kilowatt - 17-8-2008 at 19:11
I have experimented with electrolytic lead nitrating methods as outlined in that thread. I haven't tried 60Hz AC but I have tried DC and low
frequency (one square cycle every few minutes) AC. I believe the electrolysis eats up most of the ammonia, converting it to hydrogen and nitrogen
which cannot be easily recycled without the Haber Process. A lot of ammonia does escape however, and I have not yet tried to capture it in amalgam,
but that may be a good method. I have not tried any divided cells yet but have some ideas in the works for that and will continue experimenting.
Equipment cost is an issue too though and I could probably get enough nitric acid to do all the lead from the magnesium nitrate process for cheaper
than I could build such specialized cells especially with amalgam pumps and such. What I really need for an electrolytic method to work is a way to
keep the lead ions in the anolyte while allowing ammonium to pass into the catholyte. As mentioned in the lead salts thread, this is what is needed
for lead nitrate to be formed with any degree of efficiency, since the lead tends to deposit onto the cathode otherwise. The reaction only proceeds
if ammonia is removed while lead and nitrate are retained (or ammonium is at least kept separate from lead and nitrate.
I have tried lead amalgam in fused ammonium nitrate as well but it was an extremely slow reaction.
The decomposition of magnesium nitrate is not that high in my opinion; something over 300°C and I intend to do it in a RBF with a heating mantle. I
carried out the decomposition earlier at a small scale, and it was quite complete. The vapors will have to be cooled before dissolution into water,
but with excess air bubbled in the conversion to nitric acid should be complete; this is how nitric acid is made in industry from NO2. I am not aware
of any side reactions in the magnesium nitrate decomposition (for example oxygen and nitrite); it should be straight 4NO2/O2 gas coming out, as well
as water of hydration from the magnesium nitrate. Once you have nitric acid, making lead nitrate or any other metal nitrate could not possibly be
easier, as it is as simple as putting the metal into the dilute acid and letting it dissolve.
I do intend to catalytically oxidize much of the ammonia I get from the ammonium nitrate to NO2 of course, thus converting the entirety of ammonium
nitrate to nitric acid. Such a catalyst should not be required at all with the 4NO2/O2 reaction gas however. That stuff definitely forms nitric acid
with simple dissolution (better yield gotten by adding excess oxygen though).
A process which ends with glauber's salt can be considered a waste of sulfate in my opinion. In order to regain sulfate from it, high temperature
reduction with carbon as well as catalytic oxidation of SO2 to SO3 are required. Once my sulfuric acid from calcium sulfate project is further under
way (which involves exactly those two things) it would clearly be the superior method, but for now I am looking for something that does not use
sulfuric acid.
Both magnesium and lead nitrates can be considered a ready source of NO2 as well as a ready source of oxygen (the two gases are most easily separated
by condensation of N2O4), which can be quite useful to have in the lab. They are also much nicer to store than nitric acid itself. Magnesium nitrate
is a powerful oxidizer and I want to try it in some rocket propellants as I mentioned. Another fact is I am quite low on nitric acid at the moment
and need more before such a time that I can develop and electrolytic lead nitrating method.
Also for the lead refining process, the nitric acid only needs made once really. From then on the lead nitrate can be decomposed and the vapors
recycled.
[Edited on 17-8-2008 by kilowatt]Rosco Bodine - 17-8-2008 at 20:33
You can convert your glauber salt back to sodium carbonate using lime and have gypsum as the waste.
The economics of your chosen scheme are certainly related to scale, and to whatever use certain byproducts may be applied. If you are processing and
salvaging a
few dozen batteries, closing the process loop will be more expensive than simply enduring some waste by products
as the cost of doing business. Of course if you are recycling thousands of batteries as a continuing business operation, then a different rationale
applies to the investment when a net return and profit is possible for
the extra trouble of closing the loop. It may still come down to the cost of electricity and heating requirements
not being favorable economics in comparison with generating some waste byproduct or low value byproduct,
which still represents a profit or even a reduced loss for
the processing and reclamation of what would otherwise be hazardous waste. By that of course I mean that non-hazardous waste is still a profitable
outcome for having salvaged a hazardous component having some value as
reclaimed material, it is entirely acceptable to generate a
benign byproduct as waste (more or less) in the bargain.
[Edited on 17-8-2008 by Rosco Bodine]kilowatt - 17-8-2008 at 21:38
I am recycling perhaps 150-200lbs (400 moles give or take) of scrap lead which I bought years ago at recycling centers. It contains random items,
wheel weights, weather stripping, solid ingots, etc, which I melted down into ingots. It is hard lead and my main goal is to separate it into
primarily lead, tin, and antimony. Since the prices of those have gone up several times since I bought the scrap I would like to get it refined, and
I wish to have those items stocked in my periodic table in pure form, for all manners of chemistry. This isn't really for business; I just want a
method that has relatively straightforward steps, is not wasteful or expensive, and is more interesting than the more traditional methods.
Another consideration is that once the magnesium nitrate is decomposed inside a glass vessel of whatever sort, it absolutely cannot be removed. It
will be entirely cemented in and insoluble. The only options from there are to dissolve it with acid (allowing a precipitation as carbonate again but
at the cost of said acid's anions) or to react it once again directly with ammonium nitrate, using the boiling to agitate and loosen the crust. The
latter is more favorable for large volumes, but as I mentioned before the nitric acid really only needs to be generated once and the it's all up to
the lead salts. Since lead nitrate does not form highly bound hydrates I will be able to decompose the anhydrous product in a steel vessel, grinding
up the resulting litharge and smelting it. However to generate for example 250 moles of nitric acid with the magnesium method while carrying out the
precipitation each time, I would go through no less than 7 gallons of muriatic acid - unacceptable.
I have not yet done a reaction to completion of PbO and ammonium nitrate; if it is not too prohibitive that may be a good method for generating the
further nitric acid.
[Edited on 17-8-2008 by kilowatt]Rosco Bodine - 17-8-2008 at 21:42
With that quantity you shouldn't even bother trying to close the loop, just do whatever is convenient, and don't sweat the waste. If you had a
thousand times that amount, it would still be the rule, unless it was a cyclical task.kilowatt - 17-8-2008 at 21:57
What is convenient in my case is a method that produces hundreds of moles of nitric acid but does not use up hundreds of moles of sulfuric acid or
hydrochloric acid, or any other commodity for that matter. Can you imagine what they would do if I asked for 10 gallons of battery acid?
The lead is no problem since the nitric acid from it is very easily recycled, but I need enough nitric acid that I can process all the lead in just a
few batches and still have nitric acid available for other things.
[Edited on 17-8-2008 by kilowatt]Rosco Bodine - 17-8-2008 at 22:15
If you asked for ten gallons of battery electrolyte, it shouldn't even raise an eyebrow. The guy in the battery shop would go to his pallette stack
of cardboard boxed
five gallon polybagged "Qual" and take two boxes and
hand them to you. Probably he would ask you if you would like to take advantage of a special discount for a dozen or more.
I think you are multitasking this project too far to be practical ....but feel free to ingeniously prove me wrong.DubaiAmateurRocketry - 28-1-2014 at 11:51
I think salts of [Al(N2H4)3]3+ exists, However I cant find any information on it. This cation gives very high amount of hydrogen(near 10%) with
possible high density
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.