Sciencemadness Discussion Board - Aluminium Nitrate Synthesis

Aluminium Nitrate Synthesis

menchaca - 13-3-2003 at 06:00

well i have the same problem that i had before with lead nitrate i dont know if this works but..

i write this here and if it works i will be very happy

if it doesnt please tell it to me ok?
its more simple to obtain than lead nitrate
we just will need:

piece of aluminium(from a beer can aluminium paper or something like that)

sodium hidroxide

sodium or potasium nitrate

NaHSO4

the reaction is this:

2Al + 2NaOH + 2H2O ->2NaAlO2+ 3H2

NaAlO2 + 3NaNO3 + 4NaHSo4-->Al(NO3)3

+Na2SO4 + 2H2O

E.b.C.: Title

[Edited on 3-7-2005 by chemoleo]

Microtek - 13-3-2003 at 12:11

If this even works it will be extremely messy and difficult to separate. I suggest this:

Al + 3 HCl --> AlCl3 ( in water ) + 3/2 H2

Al2Cl3 + 3 NaOH --> Al(OH)3 + 3 NaCl ( do not use NaOH in excess, or Al(OH)3 will be converted to [Al(OH)4]- )

Al(OH)3 is extremely insoluble and is easily collected. Then:

Al(OH)3 + 3 HNO3 --> Al(NO3)3 ( add only just enough HNO3 to dissolve all the Al(OH)3 ).
Then you can evaporate.

BASF - 13-3-2003 at 18:41

I once made aluminum nitrate with good success using the following method:

Clean aluminum metal was dissolved in excess NaOH, lateron precipitated with the right stochiometric amount of H2SO4(aluminum hydroxide is amphoteric).
Then it was quickly pressed dry on a filter and immediately put in an erlenmeyer with concentrated(diluted acid works bad) nitric acid careful heating(not much more than 50°C) and much stirring.

The most important thing is to work the aluminum hydroxide immediately into the nitrate, because altered aluminum hydroxide has a small surface compared to the freshly precipitated hydroxide and very bad yields are the consequence.

The yielded aluminum nitrate is very hygroscopic(in the order of CaCl2 or CaNO3, for example) and cannot be made anhydrous by simple heating, because it will decompose.
Maybe with vacuum in a desiccator.

Or we use 100% HNO3 (in a stoichemetric mix instead of excess) instead of the normal 69% acid(cooling needed),
then freezing out the aluminum nitrate(boiling down under vacuum is also a possibility) and obtaining aluminum nitrate with a fair amount of HNO3 in it, but being suitable for pyrotechnic applications.

[Edited on 14-3-2003 by BASF]

BASF - 13-3-2003 at 19:02

Quote:

Al + 3 HCl --> AlCl3 ( in water ) + 3/2 H2

...this does not work.
You will get [Al(H2O)6]Cl3 instead.
Sorry if you know that already, but the way you wrote it, is misleading.

Polverone - 14-3-2003 at 11:41

One of the patents Mr. Anonymous mentioned (don't remember which, sorry) had energetic compositions that used hexamine and aluminum nitrate. The aluminum nitrate didn't need to be anhydrous, as the mixture was prepared in water and then dried to form a complex. It was interesting from a reagent preparation point of view also because it was claimed that the mixture, after ignition, yielded an especially pure and active form of Al2O3.

Microtek - 14-3-2003 at 13:22

I tried the complex from mr A's referenced patent, mixing supposed Al(NO3)3 prepared as I have stated earlier, with enough KClO4 for complete combustion of the final product. Then the proper amount of hexamine was added ( using the dry method preferred in the patent ). Nothing happened at this stage, but when a tiny amount of water was added ( less than a drop ) and the ingredients were "mashed", water was given off by the mixture and it soon turned into a slurry which dried completely without heating. This was then ground up in a mortar and pestle and testet as a solid propellant ingredient. It burned very hot but not as fast as I had hoped; it didn't compare favorably with straight KClO4/benzoate whistle mix as rocket propellant.
It may be that what I thought was Al(NO3)3 really wasn't, although my sample conformed rather well with the solubility data I have ( vs H2O, EtOH i acetone ).

PS. I have never seen the [Al(OH)2]+ species mentioned; I think all the OH- groups are neutralised on one molecule before moving on to the next so to speak. BTW [Al]3+ is common shorthand for [Al(OH2)6]3+ and the comment ( in water ) was supposed to convey this...

Rosco Bodine - 3-7-2005 at 09:26

I am uncertain if this is workable , but it may be worth experimentation as a possible method for Aluminum Nitrate .

Aluminum Sulfate is a cheap garden supply item used as a soil acidifier . What is the level of hydration is unknown .

A possible route to Aluminum Nitrate would be to form a strong hot solution of the Aluminum Sulfate and react this with
a strong solution of a cheap base like washing soda ( sodium carbonate ) or baking soda ( sodium bicarbonate ) . The insoluble unknown composition basic Aluminum salt should precipitate , but gel formation would seem a real possibility .
At a high enough temperature or a particular temperature also related to the concentration of reactants , a filterable precipitate should be possible to produce .
Baking soda would probably work better in this reaction for avoidance of a gel formation , and providing a more " acidic " neutralization condition which * may * favor the formation of an unstable basic carbonate precipitating , instead of the formation of a hydrated alumina gel as would be likely with a more basic system .

Anyway , the goal would be to not form a sodium aluminate complex , but an insoluble aluminum precipitate which contains only the oxide , hydroxide , and
perhaps carbonate constituents , with the
sodium left in solution as sodium sulfate .

The rinsed and filtered aluminum basic material could then be reacted with ammonium nitrate in hot solution , free ammonia being driven off as the byproduct . If the hexamine complex is the final product intended , then this ammonia could be utilized by being bubbled into a formaldehyde solution ,
or a water and paraformaldehyde mixture , where hexamine would be produced for later use .

Alternately , a much simplified method may work even better unless some double salt formation prevents it .
Looking at the solubilities , it seems entirely possible that double decomposition would occur for solutions of Aluminum Sulfate and Ammonium Nitrate , producing a precipitation of Aluminum Nitrate Nonahydrate and leaving the Ammonium Sulfate byproduct in solution . Just looking at the solubilities and guessing ,
about 600-650 ml of H2O
total H2O seems about right for the reaction of 1 mole Al2(SO4)3 with 6 moles of NH4NO3 , with the initial solutions mixed while hot and then cooled slowly to
0 C and filtered . The dilutions and temperatures could be optimized after a few experiments tell the story . Of course any double salt formation would just wreck this idea completely . The ammonia in the byproduct solution could be salvaged by boiling with hydrated lime , which would free the ammonia for other use .

Aluminum Nitrate has interest in nitrations
where it could be used in reaction with sulfuric acid to produce nitric acid . The concentrated warm solution of Aluminum Nitrate could be used for the purpose , and also for efficient distillations of free nitric acid from such mixtures . Since the
Aluminum Nitrate doesn't form an acid sulfate with sulfuric acid , then every 1 mole of sulfuric acid gives 2 moles of nitric acid , and the byproduct is a soluble sulfate which has likely ability to break the azeotrope nitric acid forms in aqueous systems .

Magnesium Sulfate with Ammonium Nitrate does result in a double salt .
If Aluminum Sulfate behaves the same way then this contemplated method won't work . I once looked into this scenario out of curiosity regarding the
Magnesium Nitrate for the same purposes
as would make the Aluminum Nitrate interesting . Experiments showed the
idea to be impractical . But I never tried
the Aluminum Sulfate reaction , which
it seems might work .

Update : I tried the reaction between Aluminum Sulfate and Ammonium Nitrate and got a beautiful crystalline precipitate separating from the hot solution which proved NOT to be the hoped for Aluminum Nitrate , but again a double salt or alum ,
containing both sulfate and nitrate .

So it seems to be the same story for Aluminum as for Magnesium , that the reaction of the sulfate with Ammonium Nitrate produces a double salt .

The conversion of the sulfate to a carbonate first , and then reacting with ammonium nitrate may be more worthwhile .

[Edited on 4-7-2005 by Rosco Bodine]

Zinc - 27-4-2007 at 06:02

Does anyone know how to make anhydrous Al(NO3)3?

UnintentionalChaos - 27-4-2007 at 08:38

My immediate thoughts are an alkyl nitrate added to dry aluminum hydroxide, then remove the alcohol under vacuum and store it in a very powerful dessicator

Zinc - 27-4-2007 at 13:23

Are there any other ways?

12AX7 - 27-4-2007 at 19:44

Al + 1.5 N2O5 = Al(NO3)3 ? (unbal.)

Probably done at low temperature, in Idunno, CCl4 or something.

Tim

Edit: there, happy Rosco?

[Edited on 4-28-2007 by 12AX7]

Rosco Bodine - 27-4-2007 at 20:40

That reaction doesn't balance .

Passivation would probably prevent that anyway .
Maybe an aluminum amalgam would get around that
problem . Maybe react aluminum amalgam with a molten nitrate or eutectic mixture of molten nitrates . Possibly separation in some anhydrous solvent would follow .

I can't remember now what I found about any possible double salt for Al(NO3)3 and NH4NO3 ....if there even was such a double salt . But if the Al salt follows the
same pattern as does Mg , then getting the anhydrous salt is impossible by any dehydration scheme applied to
the hydrated Al salt alone , while a double salt might be possible to be dehydrated .

I did some checking concerning the magnesium nitrate
which may be pertinent .

https://www.sciencemadness.org/talk/viewthread.php?tid=4701&...

Then the trick would be to split the double salt if you
required the aluminum salt alone .

[Edited on 27-4-2007 by Rosco Bodine]

Zinc - 28-4-2007 at 12:49

Quote:

Originally posted by BromicAcid

Quote:

Chlorine nitrate is a yellow liquid, melting point -107C., boiling point (by extrapolation) +18C. The chlorine atom may be considered to carry a partial positive charge Cl+ - ONO2-, so that in its reactions the compound is a ready source of nitrate ions. A further possible advantage of chlorine nitrate lies in the fact that its reactions are not complicated by the NO+ or NO2+ ions which are inevitably present in N2O4 or N2O5.

Because of its low melting point, reactions with chlorine nitrate, e.g.,

TiCl4 + 4ClNO3 ---> 4Cl2 + Ti(NO3)4

can be carried out conveniently at the temperature of solid carbon dioxide (-80C.) and chlorine, with excess chlorine nitrate, can be removed in vacuum at this temperature. The nitrates B(NO3)3 (-78C.), Al(NO3)3 (-7C.), and Sn(NO3)4 (-60C.) are said to be prepared in this way at the temperatures given.

Here is a method to make Al(NO3)3. I assume the anhydrous version.

Rosco Bodine - 28-4-2007 at 14:39

Looks easy as pie , so don't just make a gram or two ,
make a kilo .....
at the robotics equipped industrial facility which will be required