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ManyInterests
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testing for ammonium nitrate & nitrogen content
I've been looking to make high nitrogen ammonium nitrate to use in some nitrating mixtures, and I was looking to making ammonium nitrate via the
sodium bisulfate + ammonia + calcium nitrate, or with calcium nitrate and ammonium carbonate/bicarbonate. But I did find something a while ago
hardware stores that made me think I find some 34% ammonium nitrate.
https://www.canadiantire.ca/en/pdp/scotts-turf-builder-summe...
I looked at the SDS, they have a total of four products that are similar (all with 34-0-0 nitrogen content) but three said they were urea, the one
linked above did not say if it was ammonium nitrate or urea. They did say that when it decomposes it releases nitrogen and ammonia, this leads me to
believe that this is ammonium nitrate.
I just want to know, how can I A: test if this stuff is indeed ammonium nitrate, and B: what the nitrogen content actually is.
The second part is something I did plan on asking anyway, since if I did make some ammonium nitrate from calcium nitrate, I want to know if I got 34%
nitrogen or not.
[Edited on 4-10-2023 by ManyInterests]
[Edited on 4-10-2023 by ManyInterests]
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B(a)P
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If you are relatively certain that the material is either ammonium nitrate or urea or a mixture of both then test for urea.
Do this on a small scale as your end product will be urea nitrate, if your test material contains urea.
Make a solution of the material in water, by adding a predetermined mass of the material to enough distilled water to dissolve it.
If some solid persists, filter the solution.
Cool the solution to 5 C and place it in an ice bath.
Add, drop wise, concentrated nitric acid until a precipitate is no longer formed.
Monitor the temperature of the solution during the acid additions and keep it below 20 C.
Following the completion of acid addition take the temperature down to 0 C for a few hours.
Filter dry and weigh the solid.
You can then determine how much urea was in your initial sample, as urea nitrate is almost insoluble in the cold mother liquor.
The attached has some analytical methods for testing ammonium nitrate, including nitrogen content.
Attachment: definition-and-test-procedures-for-ammonium-nitrate_0.pdf (915kB)
This file has been downloaded 348 times
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ManyInterests
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Thanks for that. Seems like I might need to have a simpler way of testing for the time being since most of my equipment and reagents aren't with me.
I also don't have any nitric acid at the moment, but once I have some I will return to your process. Time is of the essence since hardware stores in
my area close their gardening/fertilizer sections for winter (not much market for it at this time) so time is of the essence to grab a bag. It is 4kg
so it is plenty as it is.
I was also thinking of adding some sodium carbonate to it to see if it yield sodium nitrate. But if there is urea in it, I did read that it is
possible that sodium cyanate can result?
https://chemistry.mdma.ch/hiveboard/methods/000442105.html
I know cyanates are not cyanides, but are there dangers involved with sodium cyanate?
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B(a)P
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You will not get cyanate if you do the reaction in water, or were you proposing to try without water?
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ManyInterests
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Quote: Originally posted by B(a)P  | | You will not get cyanate if you do the reaction in water, or were you proposing to try without water? |
No I was definitely not! I just simply looked up 'urea and sodium carbonate' on the internet to see what could happen. After reading that page I
linked I realized that it is not the case.
What I used to do before is add the carbonate (or bicarbonate) to the solution directly (with a large excess of water) but now I realize that it is
much more practical to just have two solutions and mix them together.
If I can also go slightly off topic, I noticed that most YouTube videos that show how to make potassium nitrate from calcium nitrate (or CAN or
ammonium nitrate) usually go with the potassium chloride route. While this obviously works, I wonder why few have made videos to show the potassium
carbonate route, since I find that route also works, and if working with calcium nitrate, the byproduct of calcium carbonate is so easily filtered out
it makes the purification process much easier. If an excess of potassium carbonate is used, then it can be cleared with hot methanol since K2CO3 is 10
times more soluble in that than KNO3.
Also it's more fun to mix the two solutions together and see the calcium carbonate crash out of solution suddenly. Maybe I should I make a video and
do some YouTube chemistry myself.
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ManyInterests
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OK, so I got the bag and finally decided to open it and get some out. It was a very rocky thing, far from what I expected. I added a large amount
(almost 500ml) of water to 77g of the stuff. It was colored nicely, but not what I expected from fertilizer prills. At any rate, I added the water and
stirred and stirred, but the pebbles did not dissolve. I did not grind the stuff, but I thought that enough came into solution (which was discolored)
since ammonium nitrate is very water soluble.
I boiled the stuff down to a low amount and dried it in the oven, it is still quite discolored, but I wasn't too worried about that. I only got around
22.5g of the stuff, which was disappointing. I carried on by measuring out 14.85 grams of sodium carbonate, which I dissolved in 300ml of water and
added the two solutions together.
It was disappointing to say the least. I was expecting fizzing and it wasn't what I got. Since I know from the stiochiometry that if it was NH4NO3
adding sodium carbonate to it will result in sodium nitrate and ammonium carbonate. I brought this solution to a boil but I did not smell any ammonia
at all, meaning whatever it is, it is not ammonium nitrate.
I will try again later with a a different method, but I have a gut feeling that whatever I got, it isn't ammonium nitrate and if it is urea, it is
more trouble than it is worth to purify.
So I do have one question (which I am going to ask elsewhere). Since I am buying a good quantity of calcium nitrate, I want to make some ammonium
nitrate from it. I know that there are multiple methods of doing it (the sodium bisulfate + ammonia that thoughtco.com published is one. The ammonium
carbonate or bicarbonate method is another (which I figured out on my own, and I am proud of that).
But while you did offer a the PDF that is a good answer... but would doing the above result in 34% nitrogen ammonium nitrate, or will it result in a
lower nitrogen %?
If it is not 34%, is there any way to increase it from whatever it is (say 27%) to 34%? I am not sure if that ultimately matters, but my main interest
in ammonium nitrate is some reactions and to make nitrating mixture for things like TNT or ETN/PETN without needing nitric acid, and to bolster WFNA
for improved RDX synth.
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Sir_Gawain
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Pure ammonium nitrate cannot have less than 35% percent nitrogen (NH4NO3= 80 g/mol, Nitrogen = 14 g/mol, each molecule of
ammonium nitrate contains two nitrogens, so it is 28/80 = 35% nitrogen). If the overall nitrogen content is less than that, it is due to impurities.
So if your "ammonium nitrate" is 27% nitrogen, it means it's actually 77% ammonium nitrate and 33% something else. You can purify it by
recrystallizing it.
“Alchemy is trying to turn things yellow; chemistry is trying to avoid things turning yellow.” -Tom deP.
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B(a)P
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| Quote: Originally posted by ManyInterests | OK, so I got the bag and finally decided to open it and get some out. It was a very rocky thing, far from what I expected. I added a large amount
(almost 500ml) of water to 77g of the stuff. It was colored nicely, but not what I expected from fertilizer prills. At any rate, I added the water and
stirred and stirred, but the pebbles did not dissolve. I did not grind the stuff, but I thought that enough came into solution (which was discolored)
since ammonium nitrate is very water soluble.
I boiled the stuff down to a low amount and dried it in the oven, it is still quite discolored, but I wasn't too worried about that. I only got around
22.5g of the stuff, which was disappointing. I carried on by measuring out 14.85 grams of sodium carbonate, which I dissolved in 300ml of water and
added the two solutions together.
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So clearly it was not 34-0-0. It seems it likely contains phosphates or maybe carbonates. Have you tried adding the 'raw' solid to a mineral acid?
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Tsjerk
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If urea nitrate is poorly soluble, couldn't you use for example sodium nitrate and hydrochloric acid to form nitric acid in situ?
[Edited on 20-10-2023 by Tsjerk]
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ManyInterests
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| Quote: | | So clearly it was not 34-0-0. It seems it likely contains phosphates or maybe carbonates. Have you tried adding the 'raw' solid to a mineral acid?
|
I definitely intend to. If it reacts and some nitrogen dioxide gets released (which I HATE to hell and back. The slightest whiff of that stuff is
nasty as all hell at best, and that's if I am lucky).
Edit: I also think that maybe trying to grind the stuff first would also help. It is designed for slow release, which makes its slowness to dissolve
something I should have expected.
| Quote: | | Pure ammonium nitrate cannot have less than 35% percent nitrogen (NH4NO3= 80 g/mol, Nitrogen = 14 g/mol, each molecule of ammonium nitrate contains
two nitrogens, so it is 28/80 = 35% nitrogen). If the overall nitrogen content is less than that, it is due to impurities. So if your "ammonium
nitrate" is 27% nitrogen, it means it's actually 77% ammonium nitrate and 33% something else. You can purify it by recrystallizing it.
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OK that is perfect. I was not aware of that.
Edit: I guess that would mean by adding ammonium carbonate/bicarbonate to calcium-ammonium nitrate it would only react with the calcium part and allow
for pure ammonium nitrate?
[Edited on 20-10-2023 by ManyInterests]
[Edited on 20-10-2023 by ManyInterests]
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Sir_Gawain
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| Quote: Originally posted by ManyInterests |
I guess that would mean by adding ammonium carbonate/bicarbonate to calcium-ammonium nitrate it would only react with the calcium part and allow for
pure ammonium nitrate? |
Yes, if you get the stoichiometry perfect. The problem is that calcium ammonium nitrate fertilizer is not always single compound with a set ratio, but
rather a mixture containing variable amounts.
“Alchemy is trying to turn things yellow; chemistry is trying to avoid things turning yellow.” -Tom deP.
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Rainwater
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I just tried my first "Kjeldahl method". I failed. Results below
| Quote: | | is the standard method against which all other methods are judged |
Thats some very strong language, it must be easy to do right. Ya. Sure lets just get some sulfuric acid, increase it boiling point, and boil the piss
out if it. Literally.
So i got some chicken @&>/ from the coup
When boiling sulfuric acid, no precautions are to extream, do it in a hole,
Far away from people, places and things
cleanup is as easy as a garden hose and a shovel.
Just a warning
by fail
i really mean that a 25ml rbf (380c) full of 18 mL 98% H2SO4, 0.500g KSO4, 0.010g CuSO4 and 0.500g of my sample kinda sorta exploded and busted a
reflux column(air cooled).
$100 bucks of virgin glass and a new thermocouple gone.
My guess is the reaction forced the nitrogen compounds to (finger quotes)
... decompose ... into nitrogen and hydrogen
"You can't do that" - challenge accepted
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Texium
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18 mL of anything in a 25 mL flask is WAY too much. Especially if gases are going to be produced. It’s no wonder it failed.
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Rainwater
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It was the only quartz i had with ground joints. It was kinda full, but after it got hot was 1/4 way up the column(14/20 200mm).
Solution had turned clear but colored green just a little.
Solids where gone.
After that point I killed the heat and started backing away. It all looked good until it popped.
Was fun, and with the proper precautions safe.
Had this happened in a fume hood, ya.
That'd be bad. Next time will be better.
Gonna use some cheap quartz tubing
Forgot to mention the literature
Following the vaguely outlined procedure in Wikipedia
https://en.m.wikipedia.org/wiki/Kjeldahl_method
TLTR:
Boil in acid, neturalize with excess base and distille into acid, evorapate the acid leaving an ammonia salt
"You can't do that" - challenge accepted
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ManyInterests
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| Quote: Originally posted by Sir_Gawain | | Quote: Originally posted by ManyInterests |
I guess that would mean by adding ammonium carbonate/bicarbonate to calcium-ammonium nitrate it would only react with the calcium part and allow for
pure ammonium nitrate? |
Yes, if you get the stoichiometry perfect. The problem is that calcium ammonium nitrate fertilizer is not always single compound with a set ratio, but
rather a mixture containing variable amounts. |
An excess is acceptable. This is because when you boil down the solution, all excess ammonium bicarbonate will decompose into ammonia gas and CO2.
When using CAN, however, the byproducts include calcium bicarbonate. This is not something that can exist in dry form (I am sure I would get an award
in chemistry if I could accomplish that feat, though) and as it is boiled down to dryness it will decompose into calcium carbonate, which is
practically insoluble in water.
Since I can get ammonium bicarbonate more cheaply and easily than ammonium carbonate (which is expensive) I will be using a lot more than that. My
idea is to use an excess of ammonium bicarbonate with either CAN or calcium nitrate, and then after boiling it down, I would redissolve. The ammonium
nitrate will all be in solution while the calcium carbonate will not. I can then simply decant/filter and dry again.
The same thing for potassium and sodium nitrate. I can get sodium carbonate dirt cheap. Potassium carbonate is a little trickier, but I can still get
a few kilos without too much trouble.
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Sulaiman
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Borosilicate glassware is ok for boiling sulphuric acid.
It is the violent bumping that scares me.
CAUTION : Hobby Chemist, not Professional or even Amateur
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ManyInterests
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| Quote: Originally posted by Rainwater |
I just tried my first "Kjeldahl method". I failed. Results below
| Quote: | | is the standard method against which all other methods are judged |
Thats some very strong language, it must be easy to do right. Ya. Sure lets just get some sulfuric acid, increase it boiling point, and boil the piss
out if it. Literally.
So i got some chicken @&>/ from the coup
When boiling sulfuric acid, no precautions are to extream, do it in a hole,
Far away from people, places and things
cleanup is as easy as a garden hose and a shovel.
Just a warning
by fail
i really mean that a 25ml rbf (380c) full of 18 mL 98% H2SO4, 0.500g KSO4, 0.010g CuSO4 and 0.500g of my sample kinda sorta exploded and busted a
reflux column(air cooled).
$100 bucks of virgin glass and a new thermocouple gone.
My guess is the reaction forced the nitrogen compounds to (finger quotes)
... decompose ... into nitrogen and hydrogen |
Ouch. Not gonna bother then. I realize that on amazon.ca they do well 25lb bags of calcium nitrate. I was wondering why they weren't available all the
time, but then I realize that, at least in Canada, the time after harvest but before the first snowfall is when farmers and gardeners spread
fertilizer over their fields. That explains why it became available all of a sudden.
No need for unnecessary risks. I got a big bag of calcium nitrate on its way and I can make all the other nitrates from it that I want/need. I will
know it is very pure since It can make 34% nitrogen ammonium nitrate.
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Rainwater
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So after a little testing with the blow torch, the glass I thought to be quartz melts at the same temperature as known borosilicate.
Gonna 1 star that seller.
I was not gentle with heating last time.
This time I heated the lead pellets up reasonably slow 10c/min. Giving water time to escape.
Really dont like the idea of molten lead, but it's the best way I know to heat this hot with accurate control.
Open to suggestions for a less toxic method?
Good news is, i dont think the ratios of acid to stuff makes much of a difference. Ran this reaction again, with gentle heating, in a 100ml rbf, same
proportions as before.
The boiling point of the sulfuric acid was around 340-345 so more K2SO4 is needed to increase the boiling point and prevent
losses.
The reaction was complete when the green color went crystal clear.
Step 1 finished.
[Edited on 24-10-2023 by Rainwater]
"You can't do that" - challenge accepted
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Sulaiman
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most heating mantles are rated for 380oC, many 450oC, some more.
Second best for me has been open-flame heating,
but temperature control is poor.
Other than handling molten metal that gives off toxic fumes and makes flasks incredibly buoyant,
what is it that you don't like about lead heating baths?
Just curious;
Does molten lead stick to glass or is it easily removed?
Do you have to remove flasks etc. from the molten lead before it solidifies on cooling?
Any other useful tips on using molten lead heating baths?
CAUTION : Hobby Chemist, not Professional or even Amateur
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Rainwater
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Flux core lead solder will wet the glass
Pure lead does not wet the flask.
Lead oxides, dross, slowly forms And can be scrapped off as it forms.
It will stick to the glass, but falls off upon cooling
From experience, alkaline compounds amd salts that are acidic will cause the lead to wet the glass
When finished, i pour the liquid lead into a pan of water to make shot pellets.
To start i place the rbf into the container and surround it in shot.
Then heat until desired temperature is reached.
To cool i lower the heating bath and turn off heating.
Im using a modified cooking hotplate with a temperature controller,
750c max, 900 watts. 30 bucks at walmart + controller
The funes are the most disgusting part, they coat everything in a toxic layer of lead oxides which is easily cleaned off with HCl
Creating an even more toxic waste.
I have used steel shot, but the downside is once the bath is lowered, you can not reinsert the appratus
"You can't do that" - challenge accepted
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RU_KLO
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@Rainwater
Have you tested sand bath?
Go SAFE, because stupidity and bad Luck exist.
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Rainwater
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Yes, the biggest problem with sand baths is the temperature gradient formed from the heat source.
Sand has a high specific heat(thermal mass) and is actually an ok insulator.
I collected some data a while back from a number of experiments
I completely insulated (top, bottom, sides) a pot about 6in in diameter and 4in tall, i was able to get the bottom to 100c while the top remained at
22c. When the bottem reached 100c, i turned off heat and waited. It took about 30 minutes for the gradient to disappear.
Sand baths are ok for long experiments, just be aware of uneven heating.
Metal shot was my next attemps, no problems with tempature gradents (<10c) but thermal transfer was limited by surface area, little balls dont have
a large contact area.
Both of these had the problem of not being able re-insert the appratus
[Edited on 24-10-2023 by Rainwater]
"You can't do that" - challenge accepted
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Rainwater
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Update:
So the H2SO4 solution from yesterday was placed into a distillation appratus whos condencer was connected into a pre weighed bubble jar containing
diluted HCl, a few drops of Methyl blue indicator was added.
Naoh was added until the solution turned from yellow to blue, then distilled until dry.
Dry weight of the bubbler jar was 162.081g
After the distillation the bubbler jar was placed a kiln at 98c until dry.
New weight is 162.083g so ya. That did not worked great.
Was hoping for more, this is not outside of the error of my scale so yield is 0%
but from a .5g of a sample, i dont know why I though i would get a measurable amount.
"You can't do that" - challenge accepted
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knowledgevschaos
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| Quote: |
You will not get cyanate if you do the reaction in water, or were you proposing to try without water? |
This sounds right to me, but the sciencemadness wiki states that urea can react with sodium carbonate in a water solvent, generating ammonia and
forming sodium cyanate. is this wrong?
I haven't been able to quote it, because it is part of a table.
I also heard that ammonia reacts with hydroxides to produce carbonates and release ammonia. Could this be used to distinguish them?
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RU_KLO
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@Rainwater just an idea, open top electric furnace?
you get high temperature control but cannot see inside....
Go SAFE, because stupidity and bad Luck exist.
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