Barium Nitrate

[s-lab]

For a while now i have been trying to find the best ways to make this good oxidiser which provides a nice green colour when used in fireworks ect....
I have thought of one way which is quite easy to do, and produces lots of barium nitrate.
First barium carbonate ( easily purchased from pottery supply stores) is decomposed to barium oxide and carbon dioxide according to this :
BaCO3 → BaO + CO2
Then the barium oxide is reacted with water to produce barium hydroxide:
BaCO3 → BaO + CO2
Finally it is reacted with ammonium nitrate (easily obtainable as a fertiliser or in instant cold packs ) to produce barium nitrate and ammonia:
Ba(OH)2. 8H2O (s) + 2NH4NO3 (s) ----> Ba(NO3)2(s) + 2NH3(aq) + 10H2O
(I hope i have balanced this right ;))

All other methods are welcome!
Thanks in advanced

[tranquillity]

What makes you think that barium hydroxide and ammonium nitrate can be made to react in the solid phase? If you did indeed mean aqueous then you would need to use hot water as the hydroxide is relatively insoluble in cold water. After that the water (and ammonia) would need to be boiled off to collect the nitrate as barium nitrate is quite soluble in water.

[s-lab]

Well hen you heat sodium hydroxide (s) with NH4NO3 (s) you will get NH3 gas being given off as NH4NO3 has a low melting point so they would react in molten state.
also as you are heating them the water and ammonia would boil off leaving solid barium nitrate.
It was just an idea and correct me if i am wrong but it seemed easy enough to me

[tranquillity]

Sorry you are correct just i think the reaction should then read:
Ba(OH)2. 8H2O (s or l) + 2NH4NO3 (l) ----> Ba(NO3)2(s) + 2NH3(g) + 10H2O(g)

I left the barium hydroxide as (s or l) as I am not sure of its melting point. Another thing to consider with a reaction involving melting NH4NO3 is that it decomposes not far above its melting point to N2O and NOx.

This reaction may also work with barium oxide:
BaO (s or l) + 2NH4NO3 (l) ----> Ba(NO3)2(s) + 2NH3(g) + H2O(g)

[Nitrato]

Pure ammonium nitrate really is a very good choice, since it will not give any metal cation that sucks the green colour (sodium,etc). But you dont need exactly to use oxide, hydroxide or carbonate.. You can just use BaCl2 in water (is my source of Barium. Technical lab grade), since the solubility of Ba(NO3)2 is extremely low, favouring it to ppt. The following data, grams of salt per 100g of water is from Seidell's book of solubility:

Ba(OH)2*8H2O : 1,67g @ 0°C ; 101,4g @ 80°C (!!!)

BaCl2*2H2O : 31.6g @ 26°C ; 58.8g @ 100°C

Ba(NO3)2 : 5g @ 0°C ; 34.2g @ 100°C

BaCO3 : 1,6 x 10^-3 g @ 8°C ; 2,2 x 10^-3g @ 18°C ; 2,4 x 10^-3g @ 24,2°C (this was take from frogfot page: http://www.frogfot.com/stuff/soltab1.png).

BaCl2 + 2 NH4NO3 ---H2O---> 2 NH4Cl + Ba(NO3)2

I used two weeks ago to make a batch of Ba(NO3)2 : a couple of grams (70g, IIRC) of BaCl2*2H2O was added in water and this was added slowly in an excess of fert. grade NH4NO3 solution (the stuff was still damp, so I didn't bothered to weight it) placed in a ice/water bath. The endotherm dissolution of ammonium nitrate helps to lower the temperature still faster. The ppt barium nitrate was filtered, the filter squezeed by hands (with gloves, of course. Most of barium salts are definately not benign, especially the more soluble salts. According with the label of my BaCl2 bag, 0.8g of BaCl2*2H2O can kill you, if swallowed/absorbed) and then the Ba(NO3)2 was recrystallized from 500mL of boiling water (as you can see Ba(NO3)2 in boiling water is less soluble than e.g. potassium chlorate, so funny). The Ba(NO3)2 are fine white cubic(?) crystals that resemble fine sugar grains and which dried with ease when spreaded in a piece of paper. Unfortunately I cant figure out how much was produced since Immediatelly after dried I used most part of it, burning with sucrose (my first Barium mix). It burns slowly with a yellow/green flame (the yellow probably is from oxidation of the carbon in sucrose..I'm going to try other mixtures to see) and gives a blackish mass of material, that I scraped of the ground and added to some HCl, filtered and then Na2CO3 solution was added, just to recover the barium values one more time .. This time the ppt wasnt that clean white stuff..Was a more impure coloured precipitate. All my effort was due the fact that where I live Barium compounds are not OTC, and somewhat expensive, so I want to recycle all that I achieve.

All the left over solutions after Ba(NO3)2 extractions/recrystalizations were combined and a solution of Na2CO3 was added, giving a white BaCO3 ppt that was decanted and washed several times with fresh water to remove at least most of sodium contamination. The BaCO3 dried and it weight a few grams (7-10g..I have no accurate scales). I was fearing that the mild acidic NH4NO3/NH4Cl/Ba(NO3)2 solution was going to just grasp on the carbonate and avoid BaCO3 formation, but when I did, the precipitation was fast and nice..Probably the Na2CO3 turned the solution basic, since I felt a slight smell of ammonia at the end.

I'm fearing that the Na2CO3 occluded some sodium inside the BaCO3 material.. So I'm now wondering If I could not use ammonium bicarbonate to make it.. Well.. There is only a good way to see =]

[Bert]

With low burning temperature organic fuels such as sugar, neither Barium nor Strontium nitrate will yield all of their Oxygen contents. You will need a higher burning temperature fuel such as Magnesium or Sulfur to reach the full potential of these as oxidizers. Also, with the Barium in particular you will need a Chlorine donor to get a good green rather than the greenish yellow you observed.

[Charles Owlen Picket]

Very good stuff Bert: That temp issue was something I have rarely heard anyone speak of. I thought that I had poor grade BaNO3 or whatever....but I have always had to get a chlorine donor and make sure that the fuel was hot burning. I don't think I even read about that previously for all the years I have been burning stuff up!

The origin of MY misconception (I believe) stems from making tracers with Gary Purrington's stuff and barium peroxide, which gave a decent green at oh dark hundred, when your eyes could fix on the streak of light. And not conceptualizing the needs for chemical colours to have "their needs" met.