I often have use for small quantities of nitrates of different metals - for pyrotechnics and for use as general laboratory chemicals. These nitrates
are easily procured by dissolving metals or their basic salts in nitric acid. Unfortunately, I don't have any real distillation apparatus, so
producing HNO3 is a very labor-intensive affair, one that I hardly wish to go through every time I need a few grams of this nitrate or that.
A few months ago I read an old message on rec.pyrotechnics that gave me an idea to try: instead of using HNO3, boil carbonates of the desired metal in
ammonium nitrate solution. The ammonium carbonate that is formed decomposes from the heat, giving off CO2 and NH3 and leaving the metal nitrate
behind.
I tried this method and found that it worked. I also found, like the original author, that the process seems to take a very long time. I had to boil
lithium carbonate with ammonium nitrate solution for more than 18 hours before I could smell no more ammonia being evolved. Far less soluble
carbonates - such as those of strontium and barium - fared much worse.
I tried to remedy these flaws by introducing hydrochloric acid to dissolve the carbonates, then adding ammonium nitrate, hoping that ammonium chloride
would be removed from the mixture (when heated to high temperatures) faster than ammonium nitrate would break down or nitric acid would be evolved. I
ended up with some heavily contaminated mixture of barium nitrate and barium chloride that nonetheless seemed pure enough for pyrotechnics. This was
the subject of a post "Beautiful Metal Nitrates" some months ago on the E&W Forum.
In reality, these nitrates weren't beautiful. They were hideously ugly - heavily contaminated by the corrosion of my stainless steel "crucible" during
the final drying stage. Only their flames were beautiful.
Alright, enough of the preamble, on to the real deal: while attempting to prepare anhydrous zinc chloride recently, it was freshly brought to my
attention that extremely concentrated solutions that remain fluid can exist at high temperatures, well above the boiling point of water. I wondered if
perhaps I might take advantage of this phenomenon - especially given ammonium nitrate's extreme solubility - to prepare metal nitrates from carbonates
without introducing any extra chemicals.
My first test was conducted with the stubborn barium carbonate: only 0.002 g of BaCO3 dissolve in 100 ml of water at 20 C, according to my references.
However, when I mixed together a stoichiometric ratio of BaCO3 and NH4NO3, placed it in a foil-lined dish on my hot plate, and added a small amount of
water, I was quickly treated to the strong scent of ammonia. I had to add more water once in a while as it boiled off and the reaction grew sluggish,
but within an hour or so it had ceased to evolve ammonia. The residue was completely dried, ground, and mixed with sulfur and aluminum powder. It
burned brilliantly and vigorously.
Then I decided it was time to try all the carbonates I had around at the time. I successfully prepared nitrates of copper, magnesium, manganese,
nickel, and strontium in addition to the barium nitrate. I am sure that lithium, sodium, and potassium carbonates would work as well. The less soluble
the starting carbonate, the harder it is to make the reaction work. The copper carbonate mix, for example, had to boil down to a very concentrated
paste before it started giving telltale traces of blue. In the case of barium nitrate, of course, I verified its nature by trying it in a pyrotechnic
mixture. In the cases of nickel and copper I was able to verify nitrate production by a change in and deepening of their colors (anhydrous copper
nitrate is one of the most beautiful chemicals I've ever seen and well worth producing just for its aesthetic merits). In the other cases I had to
mostly use my nose to detect ammonia to convince myself that there was activity.
As I've mentioned before, you may need to repeatedly apply further small amounts of water to keep the reaction going. I ran the reactions with a
slight excess of the carbonate, so that if I so desired I could produce very pure nitrates by dissolving the mixture and filtering out the small
amount of leftover carbonate.
Once you've made your nitrates it may be difficult to dry them. I was running my hotplate at about 170 C, but the strontium nitrate wouldn't give up
all of its water even after an hour of heating. The magnesium nitrate eventually formed a clear liquid (a molten hydrate, I'd guess) and just sat
there, with no apparent intention of boiling or losing further water.
All of my reactions were carried out in vessels made of or lined with aluminum foil. It was unharmed even by copper and nickel nitrate, and it was
easy to obtain the dried nitrates from it by bending the foil so that the solid crumbled away or could be peeled out.
One final note: I was initially nervous about heating ammonium nitrate and copper carbonate since NH4NO3 + Cu = DON'T! when it comes to explosive
safety. However, I had 5 grams or so dry completely without incident. When I placed a few bits of the CuNO3 directly on the hot plate surface, they
just dried, and didn't react at all until I added a bit of sugar, which caused a small spurt of flame after a short delay. Keep in mind that this was
all around 170 C and I can't vouch for stability at higher temperatures.madscientist - 22-8-2002 at 16:50
Wouldn't the easiest route to, say, Ca(NO3)2, be reacting Ca(OH)2 directly with NH4NO3?
Ca(OH)2 + 2NH4NO3 --> Ca(NO3)2 + 2NH3 + 2H2O
By the way, I use pH paper to detect ammonia, rather than my nose.
Hydroxides
Polverone - 22-8-2002 at 17:50
Hydroxides will work as well. In many cases it's easier to obtain the carbonates of metals rather than the hydroxides, and I knew that carbonates
would be the more challenging test, so that's what I worked with. For me, it's as easy to buy calcium nitrate as it is to buy calcium hydroxide or
ammonium nitrate. I'm sure everybody's circumstances will differ, though. You can use pH paper for detecting ammonia but the nose is at least as
sensitive and completely reuseable. Save the planet and all that.rikkitikkitavi - 25-8-2002 at 01:15
One possible explanation:
NH4+ is a stronger acid than H2CO3
Thus
2NH4+ + CO3-- => CO2+ H2O + NH3
you are just using a little water, solubility of NH4NO3 is extremely high @ 100 C so the concentration of NH4+ is high,.
(the temperature is probably a lot higher due to dissolved NH4NO3)
That and the fact that CO2 and NH3 doesnt show much solubility in water @ 100 C shifts the reaction to the right.
Left is NO3- ions and M2+
/rickardIodineForLunch - 31-8-2002 at 14:27
Note that this method can also be used to produce metal perchlorates by substituting ammonium perchlorate for the ammonium nitrate.
Where could one purchase lithium carbonate? None of my pyro chemical suppliers carry it.
David HansenIodineForLunch - 31-8-2002 at 14:50
I'm going to list the dehydration tempatures of a few metal nitrate hydrates, for convenience's sake.
Mg(NO3)2.3H2O - 100 degrees C
Mg(NO3)2.6H20 - 330 degrees C
Sr(NO3)2.4H20 - 100 degrees C
These were referenced from the CRC Handbook of Chem & Phys, of which I am lucky enough to own a copy.
I couldn't find the information on copper nitrate or manganese nitrate.
BTW, was your nickel nitrate green, or white? If it was green, it's the hexahydrate form. If it's white, it's anhydrous.
Could you describe the aesthetic properties of anhydrous copper nitrate in detail, please?
David Hansen
An artwork ....
BrAiNFeVeR - 1-9-2002 at 11:35
If it is what I think it is, it comes in beautifully large dark blue/turquise crystals (hexagons iirc ? )
I once made it when dissolving some copper in 65% HNO3, and the next day the entire bottom of the beaker was full of it !!
I immediatly went to show it to my mom, knowing that it wouldn't last long in my humid-as-hell-climate ...
Even when left in a dry warm place, it just sucks up all moist it can find :-(
Nitrate fun... continued!
Polverone - 1-9-2002 at 13:27
I obtained my lithium carbonate as a powder intended for making ceramic glazes. Find a ceramics supplier. They have all sorts of useful chemicals at
very low prices. Visit www.clayartcenter.com and look at their dry chemicals list to get an idea of what sorts of things you may find at ye olde ceramics supplier. I
have, myself, obtained manganese dioxide, nickel carbonate, copper carbonate/hydroxide, lithium carbonate, strontium carbonate, barium carbonate, tin
oxide, cupric oxide, ferric oxide, and chromium oxide from a ceramics supplier. They're also a good source for exotic transition metals like vanadium
and cobalt.
Unfortunately, lithium compounds are just a curiosity for pyrotechnics unless you live in a dessicator. When I made my lithium nitrate I dried it in
the oven and immediately ground several grams in my mortar and mixed it with sugar to test it. It burned vigorously, and with a beautiful reddish
flame, but the leftover powder was damp to the touch within 10 minutes, and actually wet in 30 minutes. This was on a sunny day in the middle of the
summer. It would be far worse in a damp climate. I haven't yet tried mixing lithium carbonate (non-hygroscopic) with other oxidizers to get the color
without the moisture affinity.
The nickel nitrate I made was green; I know it wasn't anhydrous. I never tried to fully dry it.
The copper nitrate (anhydrous) that I made (think I made - copper compounds are sneaky) was a striking deep blue with a purplish cast to it. In a
second test where I just added copper carbonate/hydroxide (small amounts) to molten ammonium nitrate, the fluid turned a deep blue that slowly
lightened over a period of days (stored in open air).Xenos - 1-9-2002 at 15:26
I would like to get some lead nitrate. Is that sold under a brand name, or would i have to go about the method to make it? Or more importantly, could
it be made with that method, and to what extent of difficulty?IodineForLunch - 1-9-2002 at 18:32
Xenos, that would depend on the availability of lead carboante. :-/
David Hansen
Lead nitrate: the hard way
Polverone - 2-9-2002 at 21:50
If you boil lead metal (preferrably filings or another form with high surface area) you will smell ammonia and get a white powder (as well as
miscellaneous other stuff including, I think, lead nitrite). I'm not sure exactly what this powder is, but it is a basic lead compound that does not
dissolve in water. It can, however, be dissolved in acetic acid to form lead acetate. If you so desired it could then be reacted with potassium/sodium
carbonate to form insoluble lead carbonate.
You might be able to react this with ammonium nitrate to form lead nitrate. A tedious and roundabout method, I know. It's only worthwhile if you
really can't obtain nitric acid. I haven't even checked (yet) to see if the carbonate reacts with ammonium nitrate in a fashion similar to the other
metals I've tried. I've reacted the lead acetate solution with calcium hypochlorite to form lead dioxide.madscientist - 3-9-2002 at 13:07
The white powder is probably Pb(OH)2.PHILOU Zrealone - 8-9-2002 at 10:28
Metal nitrates can be done from NH4NO3 and Metal oxydes, Metal hydroxydes, Metal crabonates.
So basically:
2NH4NO3 (s) + Na2CO3 (s) -heat-> 2NaNO3 + (NH4)2CO3
As you all know:
(NH4)2CO3 is chemical yeast and upon standing (especially in hot environement) free NH3, H2O and CO2.
So all gaseous and low boiling point liquid will leave the system in equilibrium and let the non volatile nitrate!
This procedure can be done in solution with saturated NH4NO3 (it is indeed safer no to heat dry NH4NO3 especially with another compound).
The same process applies to Hydroxydes since NH4OH is unstable and volatile and free NH3 and H2O.It also works with oxydes that are the anhydride of
hydroxydes after all!
2NH4NO3 + Ca(OH)2 -heat-> 2NH4OH + Ca(NO3)2
NH4OH --> NH3 + H2O
2NH4NO3 + CaO --> Ca(NO3)2 + NH3 + NH4OH
Beware of NH3 (lacrymator, eyes destructor, war gas- makes your lung bleed on long exposure).
Also I would recomand to work with tiny amounts and with excess water when working with complexating metals like Cu, Fe, Ni, Co, ... because you all
know some of those metals do catalyse the decomposition of NH4NO3...maybe via some explosive polyamino nitrate complexes (TACuN, TANiNn, TACoN are
three that comes to my mind) but also Cr aminonitrate complexes ....
So the procedure is quite safe for alkalino metals and alkalinoheart metals but
beware of transition complexating metals !
PH Z karandikarmv - 4-9-2003 at 22:10
i am looking for caclium nitrite from calcium nitrate... how to proceed... cause caclium nitrate.tetrahydrate starts decomposing to CaO at ~580-620
degC...
i have tried several ways but could not Succeed... please let me know
Lithium Nitrate
chloric1 - 1-3-2005 at 13:25
Found an intersting patent (US 6,555,078) and proceeded armed with sodium nitrate,lithium carbonate, and muriatic acid. I neutralized the acid with
the lithium carbonate and dried until I had moist chloride crystals that where acid free. I then dissolved these in water and dissolved a 20% excess
of Sodium Nitrate in near boiling DI water. I mixed everything and concentrated the solution until cloudy. Filtered out NaCl and let cool. Some what
appears to be NaNO3 crystalized and I filtered again. Added 100 ml of acetone and some more solids deposited. I filtered to get a crystal clear
solution and boiled all acetone and some water went with it. I have a few grams of inorganic crust at the bottom after coooling with still ample
solution left to test. I added methanol and I wonder how the solubility of the lithium nitrate is in methanol? Anyone have the right books? My
books(Lange's, CRC 1975) does not inform me enough.
I do know that sodium nitrate is only soluble in methanol 1 part in 300 parts. THe alcohol is about -3C and I broke up the crystaline cake and it got
cloudy. Good sign I think I decided to turn on the heat. Might agitate too.
[Edited on 3/1/2005 by chloric1]
[Edited on 3/1/2005 by chloric1]
Failure!
chloric1 - 1-3-2005 at 15:54
Well the methanol solution did not take place and then I decided to test by mixing with sugar and igniting. Just yellow sodium flames
I could of done a sodium carbonate test on the solution but if the resulting salt cannot even produce a lithium red than it does not matter how much
lithium I had in my nitrate. Mind as well been pure sodium Nitrate.
Next time I will just heat the lithium carbonate with ammonium nitrate. I guess lithium bromide mind as well be written off my accessability list for
now.neutrino - 1-3-2005 at 19:32
According to Merck, LiNO<sub>3</sub> is ‘soluble in alcohol’.chloric1 - 2-3-2005 at 10:29
Yeh that was about the extent of the information I could get. Some of the material went into solution in methanol but I guess it was less than 5%. My
lithium carbonate is from a ceramic source and I need to verifiy it contents some other fashion.Quince - 6-3-2005 at 02:00
The only potassium available OTC here is the sulfate (and it's very cheap). Is there any way to transform K2SO4 to something useful, such as a
nitrate? When I search, it's almost always on the right side of an equation, but I've seen the following two reactions:
K2SO4 + Ca(OH)2 --> 2 KOH + CaSO4 (German site)
K2SO4 + Ca(NO3)2 --> 2 KNO3 + CaSO4 (roguesci)
Will these reactions actually happen?
I also have NaOH and 12% HNO3 as reagents, but they cost a good deal more than the lime and K2SO4.
I also read that when heated, K2SO4 produces SO2 vapours, but what's left, K2O? Can that be transformed into a useful salt?
What about electrolysis?
[Edited on 6-3-2005 by Quince]neutrino - 6-3-2005 at 06:54
The second reaction should work, but the first one would be a pain as the Ca(OH)<sub>2</sub> is only slightly soluble in water. This means
that you’d mix dilute solutions, filter out the precipitate, and boil the stuff down a lot. A word of warning: calcium sulfate precipitates are very
hard to filter, as they are more gelatinous than powdery.
Heating would probably give you some oxide, but the temperatures required to get any yield would make it impractical. Judging by what I’ve read on
these forums about people’s experience with electrolysis, it would give you a small amount of product, but nothing worth recovering.Quince - 6-3-2005 at 07:19
Well, I make Ca(NO3)2 by neutralizing dilute HNO3 with Ca(OH)2. It's hard to get all the lime to react as the sulfate forms a coating on the
undissolved particles.
Since I have the first mix in boiling water already, but it doesn't seem to react (much of the lime still falls to the bottom), I'll try
adding the HNO3 to the mix when the water evaporates.
What about using NaOH or NaNO3 as a substitute for the calcium ones in these equations?
[Edited on 6-3-2005 by Quince]neutrino - 6-3-2005 at 09:02
Sorry, sodium sulfate is very soluble.Quince - 6-3-2005 at 20:59
Adding HNO3, and with a lot of mixing and some heating, it worked. I put it in the freezer after filtering and the KNO3 crystallized. Yield was
crap, about half, but until I can find somewhere to buy it, I'm stuck either with this or NaNO3.JohnWW - 7-3-2005 at 13:31
Some posters on this thread seem to forget that practically all metal and ammonium (including substituted) nitrates are soluble in water, and in many
other polar solvents as well. The major exception to nitrate solubility is covalent nitrate esters of high molecular weights.
In addition, the nitrates of lithium and of polyvalent metals (except barium and radium, the most electropositive ones) can have water of
crystallization, in addition to anhydrous crystals, depending on temperature and pH of crystallization.Quince - 7-3-2005 at 23:06
Guessing that by "Some posters" you mean myself (I can't see who else), I don't understand what it is I posted to call for your
comment. Where did I claim they are not soluble?AngelEyes - 2-5-2006 at 06:43
Going to resurrect an earlier thread rather than start a new one.
I have some Copper Nitrate and Barium Chloride. I am trying to precipitate Barium Nitrate from an aqueous solution of the two.
Barium Nitrate (8.2g/100ml water @ 20c) has a much lower solubility in water than Copper II chloride (42.6g/100ml water @ 20c) so I am expecting it to
precipitate out. Is this feasible or am I way off track here?
I have used a stoichiometirc ratio of materials.
I want the Barium Nitrate for use in green lances and portfires, the BaCl2 just doesn't cut it. Also, the Copper Nitrate is green, I am assuming it's
a tri-hydrate - .3H<sub>2</sub>O?
Thanks
AngelEyes.woelen - 2-5-2006 at 10:45
If your copper nitrate is green, then it is not pure. I also have copper nitrate, and it is a beautiful crystalline glittering solid, bright blue,
even nicer blue than copper sulfate. Its solutions also are nice blue.
Don't you have any other nitrates around? If you have ammonium nitrate from fertilizer (possibly mixed with chalk), then you can use that. Just
dissolve in as little as possible of water and let stand for a day. All crap settles and then you can mix that very concentrated solution of ammonium
nitrate with your barium chloride solution.
KNO3 also may do the job, but it is somewhat less soluble. NaNO3 is very soluble, but I do not recommend that, because your product will always
contain a little amount of the sodium ions. These produce very strong orange/yellow light and certainly will spoil your compositions. Traces of
potassium ions or ammonium ions, however, do not really harm.
Actually...
AngelEyes - 3-5-2006 at 01:24
...my Copper Nitrate <i>is</i> blue. Not sure why I put it as green. It's large, clumpy blue crystals. Looks like blue quartz.
I do have Sodium Nitrate but don't want the pollution from Na. I have no other Nitrates, but I figured this reaction would proceed OK with the
materials I have.
I have some precipitate drying, it has a slightly greenish tinge though this could be contamination with Copper Chloride I s'pose. I'll see how it
dries then try some flame tests to try and determine what I have.
Cheers
AngelEyes.AngelEyes - 11-5-2006 at 09:42
Right...so I have a quantity of a light-turquoise powder after filtration. It's taken ages to dry but seems non-hygroscopic now that it is dry.
If I try a standard flame test on it, it just melts and turns black. There's a tinge of greeny-blue...but that's hardly conclusive. I guess I must
have a mix of compounds in there or something.
If I mix with sulphur and try to ignite it, it turns black and I can see the dance of the light blue sulphur flame on top...then it eventually ignites
with a bright white, plus a hint of green, flame. It burns slow and leaves a black residue. Maybe sulphur isn't the best fuel to use here.
One odd thing is that when the powder was not quite dry and I mixed in a little hexamine it turned a sandy beige colour before my very eyes. Still
didn't burn well though...neutrino - 11-5-2006 at 15:35
Hexamine is an amine with lone pairs. The Cu<sup>2+</sup> ions probably got bound to these.Misanthropy - 31-5-2006 at 17:06
Quote:
Originally posted by IodineForLunch
Where could one purchase lithium carbonate? None of my pyro chemical suppliers carry it.
...in a Calcium multi action fertiliser. Seems it's a mix of Ammonium Nitrate (about 80%) and chalk or other calcium compound. It's assay says it
contains both nitric and ammoniacal nitrogen, as well as a soluble calcium compound.
I added about 500ml of water to the 750g contents which I figure will dissolve all the nitrate nicely. It does, and I am left with some undissolved
crap. I filter this off to leave a nice, clear solution of what must be mostly Ammonium Nitrate. I add some Ammonium Sulphate to precipatate any
Calcium Nitrate out but there's no reaction so I figure I have quite pure Ammonium Nitrate solution with no dissolved calcium ions. I then add some
concentrated Barrium Chloride solution and get a precipitate fairly soon. I leave this for a while, then come back and filter off what I presume must
be fairly pure Barium Nitrate.
However, after drying and trying to ignite with various fuels I get hardly any reaction at all.
Do you reckon I have a mix of Barium Nitrate and Ammonium Chloride rather than mostly just the barium nitrate?
Have I got my equations / solubilities wrong?
Is Barium Nitrate just a real PITA to make by this method?
Should I redissolve (or re-wash) the solid I have?
I can do a good red fire, as well as yellow, white, purple-ish and orange. Blue is hit and miss anyway, but green <i>should</i> be quite
straightforward dammit! I use Strontium Carbonate for the reds and it's really good - is Barium Carbonate equally good for greens or can you really
use only the nitrate or (per)chlorate?
Beofre anyone mentions Skylighter <i>et al</i>, be aware that I am in the UK and we have a nanny state. Plus, I am quite skint and so
would prefer to use what I currently have.
Cheers for any pointers
AngelEyes.12AX7 - 9-6-2006 at 11:32
I have good results making barium nitrate in a similar way (from a less concentrated fertilizer also containing potassium sulfate and calcium
phosphate). The sulfate is precipitated with Ba and washed, then excess Ba is added to form Ba(NO3)2 on evaporation. It forms regular, coarse
crystals with a hopper tendency.
Er.. um... hummm... did you happen to think that your ammonium sulfate would be causing the barium to precipitate!?
My product burns in starts with magnalium, being a basic strobe comp. Bright greenish light.
TimAngelEyes - 12-6-2006 at 03:23
Quote:
Er.. um... hummm... did you happen to think that your ammonium sulfate would be causing the barium to precipitate!?
Well, no - I added the Ammonium Sulphate in the first step to precipitate out any Calcium ions that may have been there -, but now that you mention it
any soluble suplahte <i>would</i> ppt out any Barium in solution wouldn't it? But I added the sulphate before any Barium Chloride goes
near the solution, got no ppt so assumed that all the Calcium was the insoluble crap I filtered off.
I hav another batch drying that used excess Ammonium Nitrate solution. I will dry the hell out of it then try again with other fuels...like sugar
maybe.
Cheers
AngelEyes
[Edited on 12-6-2006 by AngelEyes]
strontium, manganese, lithium nitrates
tetrahedron - 14-11-2012 at 15:23
i had some success isolating strontium nitrate from strontium carbonate and ammonium nitrate, but the procedure is slow. see my post here.
with manganese it seems to proceed analogously, but i wasn't able to isolate the manganese nitrate due to it's good solubility.
i'm currently making lithium nitrate. the carbonate has some solubility, and this reaction appears to work much faster, and the ammonia evolution to
be much stronger. i improvised a distillation setup and am now recovering quite concentrated ammonia solution. more details to follow.
<sub>2</sub> + BaCl2 --> CuCl2 +
Ba(NO<sub>3</sub>