What are metal nitrates?

Answer 1

Metal nitrates - the easy way

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.

Answer 2

Metal nitrates are compounds formed by the combination of a metal cation with the nitrate anion (NO3-). They are commonly used in various industries, such as agriculture for fertilizers, in pyrotechnics for producing colorful flames, and in the production of metal oxides.