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Murda
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Synthesizing UN with calcium nitrate
Well I new to the forum but not new the high explosives. The synthesis I use for urea nitrate is:
All that's needed here is urea, ammonium nitrate, and some HCl. Two solutions are made - 39.5g of ammonium nitrate in 25 ml water, and 27g urea in 35
ml water (scale up as needed). These two solutions are then mixed, leaving a clear solution of the two. This solution is then heated to 80 degrees C
in a water bath. Once the temperature is reached, the beaker is removed from the water bath. Then, 50 ml of 31.45% HCl is added to the solution with
rapid stirring and left to steam. This is then allowed to cool with further stirring. Crystals of urea nitrate will become visible in the solution
after some cooling has occurred, and stirring is continued. The solution is allowed to cool with further stirring, breaking up the clump of crystals
every once in a while. The solution is then placed in the freezer for a few hours, or overnight, and the clump of crystals once again broken up. The
crystals are then filtered and left to dry.
I was doing some thinking. Since AN is kind of hard to come by maybe substitute the AN with CN using about 75g of CN and everything else stays the
same. But I did some reading in the MSDS of prill CN and the melting point of CN is 45C. Will this pose a problem when heating the solutions to 80C?
Also I read somewhere that CN can emit toxic fumes. I can't seem to find it anymore. Let me know your thoughts and I will try the synth tomorrow if
everything looks good.
Thanks
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chief
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Well: I never make any explosives (much better sound comes from my electric guitar  , and I need the fingers for that, as well as the eardrums ...  ),
but:
==> Is the above post bullshit ? Or does it work differently from standard nitration ? Because: So much water involved ...
If it works: Is it good for any other nitrations ?
As I said: Explosives are not my stregth ...
===============
Or is urea-nitrate just some sort of a "salt" (and not a nitro-explosive) ? Then it would be a forced crystallization: NH4NO3 anyhow has a _very_ high
molar solubility ... and the addition of HCl would ppt. the compounds ...
With Ca(NO3)2 it would then be the question, if the ratio of its solubility to that of the chloride would be similar to the ammonium-compounds ... :
Then it also might work similarly ...
[Edited on 6-3-2010 by chief]
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hissingnoise
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If tomorrow's synth fails and it might, you'll have to make AN from NH4SO4 and Ca(NO3)2 by metathesis.
It's extra hassle but at least CaSO4 precipitates!
Or if you have H2SO4 you could go straight for HNO3. . .
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Rosco Bodine
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The reaction of hot strong solutions urea and ammonium nitrate and HCl to precipitate urea nitrate is driven by the solubility relatiionships of the
reactants and products, and variations upon that process should also work unless the complication arises involving formation of double salts and/or
hydrated salts
having solubilities which compete with the desired precipitation of urea nitrate,
in which case a coprecipitate may result instead. Reactions involving substituted potential alternate nitrates and/or acids are worth trying and only
experiment will
reveal what works and what doesn't. The only variation which I have tried is
the substitution of H2SO4 for the HCl in the original scheme, and that did appear to lead to a double salt or contaminated precipitate, however there
was no further investigation or even confirmation of that result by me. The concentrations of reacting solutions and the temperatures may have to be
adjusted and optimized to different figures better suited for variations on the
original process which I shared here in this forum several years ago.
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Murda
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Thanks a lot for the information guys.
@hissingnoise Well I have AN and have gotten the synth to work and produced great yields using AN but It's just a huge hassle to get small amounts.
That's why I am trying to tweak the synth a little with a less know nitrating salt which I can get in large amounts. CN around here is quite expensive
but it's better than nothing. The reason I am using this synth is because nitric is impossible to come by that is concentrated unless of course I make
it myself and yet too valuable for UN. The reason I am tweaking the synth is because I would like to use AN for synths like ETN which are more
powerful than UN. It also got me thinking if this synth works with replacing CN maybe it could also be possible with ETN. The only problem I see is
that in the CN MSDS it lists H2SO4 as an imcompatibility. Not sure if it would be the same with a CN solution. I bought CN because I think it has
great potential. For instance in ANFO mixtures. In this patent CNAN/FO A mixture of 60% AN 40% CN 5.5% hydrocarbon makes a larger diameter than ANFO by itself by over a foot.
I also feel that maybe a mixture like CN/soy might work as well. I just have no where to do my testing since the area I used to test was replaced with
some government factory. I have heard CN/NM is stronger than AN/NM. MY CN is from the fertilizer CO yara and they clearly use their CN in their
explosive mixtures as seen here. Yara explosives
Someone willing to do some testing? It would make me very happy. I shall try the UN synth today. If I don't return I guess you know what happened...
By the way the melting point of CN being 45C and me heating the urea/CN solution to 80C won't pose any dangers?
Thanks
[Edited on 3-6-2010 by Murda]
[Edited on 3-6-2010 by Murda]
[Edited on 3-6-2010 by Murda]
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not_important
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That's not truly the melting point of the calcium nitrate, the CN is a hydrate and what is happening is that it is dissolving in it's own water of
crystalllisation about 45 C. Obviously if you were using equipment to handle solid prills of CN, its converting to a flowing solution might cause
problems, if it's already in solution there's no change at 45 C.
The somewhat vague amount of water in the solid CN will make it difficult to measure exact amounts of Ca(NO3)2 needed for use with X grams of urea.
You'll need to use enought to be sure to have some excess CN.
Both CaCl2 and Ca(NO3)2 are quite soluble, there should be no problem from that standpoint. Calcium salts form weak complexes with many ammonia
derivatives (and water derivatives the alcohols), there are lists of substances that shouldn't be dried with CaCl2 because of this. If urea is one of
those there might be yield or contamination problems.
Calcium nitrate is 'incompatible' with H2SO4 because it forms nitric acid and nitrogen oxides as a result of the heat released - the H2SO4 sucking up
all the water in the CN prills; nitrogen oxides are not good to inhale.
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Murda
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The CN is about 1/16" prills that will be added to water and heated with the urea solution.
EDIT: So I did the synth and it was different than UN. I'm not sure what I just made but after pulling it out of the hot water bath I noticed very
tiny crystals it seemed like and I added the HCL and started stirring rapidly. It took longer than UN but it started forming crystals before my eyes.
Only thing is it wasnt the big UN crystals like I have had experience before. They were very tiny like AP. after letting it sit for about 5 minutes I
noticed tiny crystals floating just like AP too! But to me it looks like UN to me. It's now sitting in the freezer. By the way I doubled the amount of
water and CN but kept everything else the same.
Thoughts? I will keep you updated on if anymore crystals form and have a video as well right after the crystals started forming. I might put it on
rapidshare if people want to see it.
EDIT: More crystals have formed. Still tiny and no clumps either.
Thanks
[Edited on 3-6-2010 by Murda]
[Edited on 3-6-2010 by Murda]
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Melgar
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Here's what I think happened: NH4Cl has a lower solubility than UN, so it precipitates out first. CaCl2 is so water-soluble it's exothermically
hygroscopic, so I certainly wouldn't expect it to precipitate out any time soon. And UN has a very high solubility, so expect difficulty
precipitating that out too. However, calcium nitrate fertilizer is a double-salt hydrate with ammonium nitrate and water, so I think those little
crystals you're seeing are ammonium chloride crystals precipitating due to the presence of ammonium nitrate in the original salt. Sulfuric acid would
precipitate out the calcium sulfate and probably force the formation of urea nitrate. Or urea sulfate, but probably urea nitrate. That's a lot of
salts to be accounting for at once.
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Rosco Bodine
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Why would you think that NH4Cl has a lower solubility than urea nitrate ?
If that were the case then the original synthesis would not proceed in the way that it does, where urea nitrate precipitates from the byproduct
solution remaining
and that byproduct is predominately NH4Cl. Are we doing real chemistry here
using stoichiometry and solubility and temperature changes in a deliberate way
hoped to likely produce an intended result ? Can you write the formula for what it is that you are trying to do ?
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Murda
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I was simply trying to replace AN with CN since CN consists of AN to see what happened. There was no real chemistry here. I would however like to
discuss the many hidden uses of CN like I mentioned above. Not sure if it's ok to make another topic on this subject though. I think it has great
potential and is not a well know explosive.
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Rosco Bodine
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You need to have an understanding of the formulas and weights and the solubilities of the materials you are reacting in order to set up your
experiment
and make any sense of the result. If you are substituting a hydrated double salt of Calcium Nitrate - Ammonium Nitrate you should do the math with
regards to
label on the bag which should state what weight percentage of the contained Nitrogen is derived from nitrate. It is a requirement that fertilizers
state what percentage of the component nutrients that the fertilizer contains. The common
hydrated double salt of CN - AN is probably labeled 15.5 - 0.0 - 0.0 analysis,
based on a formula 5 Ca(NO3)2 - NH4(NO3) - 10 H2O , with 14.4% of the Nitrogen
being "derived from nitrate" and the other 1% "derived from ammonia".
Therefore, per kilogram of the fertilizer there will be 144 grams of Nitrogen derived from nitrate, and the mole weight of Nitrogen is 14.0067, so
there
is about 10.28 moles of Nitrogen (and nitrate as NO3) per kilogram of that fertilizer.
The mole weight of urea is 60.06, so about 10.28 X 60.06 grams, 617.4 grams
of urea will be required for reacting with each kilogram of the fertilizer.
The common 31.45% HCl is almost exactly 100 ml per mole and 10.28 moles
or 1028 ml of HCl will therefore be needed.
A kilogram of the fertilizer should dissolve in 450 ml of H2O at room temperature,
and less H2O would be required for hot solutions.
These figures are unverified, as I have not checked the analysis percentages
stated for the fertilizer against the hydration to see if the figures reconcile,
but these numbers should give you a better idea of how to proceed with an
experiment using some ball park numbers.
I just crunched the numbers a bit on this and the stated numbers are pretty close if not exact. A mole weight
for the fertilizer works out to about 1080.49 grams which
contains 11 moles of nitrate expressed as NO3. For that mole of fertilizer as starting material, you would need 1100 ml of the 31.45% HCl and 660.7
grams of urea.
Adjusting proportions for a reaction neutralized by 1 liter of the HCl which is a convenient reference reagent, 982.3 grams
of the fertilizer and 600.6 grams of urea would be required
to balance the reaction. The optimum water content and crystallization procedure and filtering temperature is something which will have to be
determined by experiment.
In terms of the conversion equivalence used for substitution
of the hydrated CN - AN double salt, used in the place of AN,
the multiplier is 1.2272 X the AN amount as would balance
the same reaction equation.
Always, you should first do the math ....
then do the experiment. You must do the math first
in order to have a practical knowledge of how much of each reactant is required by theory if the reaction goes to completion for the balanced equation
for the reaction,
and to calculate percentage yield for your actual product.
It is those numbers which help you to interpret observations of the experiment and if necessary to adjust the process.
The optimum yield based upon urea will likely not be achieved using a process based on proportions balanced exactly according to theory, because there
is a dynamic involving the solubility and pH which affects crystallization
of the desired product. So the balanced equation for any reaction is only a guide and a practical synthesis usually requires some trial and error
varying the proportions to
determine what changes tend to drive the desired reaction further towards completion and/or produce the product desired in a form and purity and
quantity which is a
compromise against the 100% of theoretical yield that
is only sometimes achieved.
If I recall correctly, because urea nitrate is very acidic, its precipitation is favored by a reaction mixture having an excess of acid, which was
used at 11% in excess of theory
based on the urea, while the nitrate was used at about 9.7%
in excess of theory based on urea. If the same adjustments
are favorable for the reaction substituting the hydrated double salt of CN - AN, then the parallel synthesis adjusted
accordingly away from the precisely balanced proportions
would be as follows:
For the reaction referenced to 1 liter of 31.45% HCl being added to neutralize and drive the reaction by excess acidity,
540 grams of urea in 700 ml H2O ( 9 moles urea )
970 grams of hydrated double salt CN - AN fertilizer in 500 ml H2O ( 9.9 moles NO3 )
1000 ml 31.45% HCl ( 10 moles HCl )
At an 84.4% yield which would be the same as for the original process using AN fertilizer, 935 grams of dried UN will result.
Due to the high solubility of the anticipated CaCl2 byproduct,
it may be possible to use even more concentrated solutions
for the synthesis and obtain a higher yield of urea nitrate than by the method using AN as the nitrate. However,
that outcome is uncertain and predicated upon the absence
of any complication involving formation of any competing double salts.
The maximum theoretical yield of 100% based upon urea would be 9 moles of urea nitrate, 9 X 123 = 1107 grams
In the original poster's proposed reaction, the quantities
are correct with the exception of the substitution of 75 grams
of the CN - AN which should be 48.5 grams instead.
[Edited on 7-3-2010 by Rosco Bodine]
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Murda
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Very awesome. Thanks a lot. I'll replace the 75g of CN with the 48.5 and tell you what I get. Whatever I made before had some oxidizing powers. Of
course from the CN a 50/50 ratio ?/AL burned like a flare but needed substantial heat to start it. A little sprinkle of AL on the top of the mixture
did the job.
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Murda
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SUCCESS!!! I have a video I would like to share with you guys. The crystals aren't as big as when I tried it with AN but it definitely looks like UN.
I did however do bigger batches before with the AN but this one was smaller for testing. I will try to filter and dry the UN tomorrow but I will be
busy all day and might have to wait the day after. When the UN is dried I will make another video. I am unable to det the UN because I no longer have
a "spot". Some government plant is in it's place now. I want to try and substitute AN with CN to synthesize ETN but do not want to be spoon fed. Where
can I do some reading on how to figure out the ratio of CN is needed? Here is the synth.
For this experiment I used 5 grams erythritol, 20 grams ammonium nitrate, and 30 ml sulfuric acid (to allow for ~35% excess nitric acid). This can be
scaled up.
The ammonium nitrate was added to the sulfuric acid and allowed to dissolve. This was then chilled in a fridge. Once the acids were cooled the
erythritol was added in portions with stirring, while keeping the temperature down. Temperature control is generally more of an issue when using
larger amounts than this. The mix was stirred for a few minutes after the final addition and then left to sit for about half an hour longer with
occasional stirring. This was then poured into about 500 ml cold water and a precipitate was instantly seen to settle to the bottom. After washings
with water and some sodium bicarbonate in water, the product was left to dry. 8g is the yield of ETN.
When I figure it out I will report back to you to see if it's right. Only thing I had to change was I had to add an additional 5 ML of water to fully
dissolve the CN because no matter how much I mixed it the prills were not getting any smaller. I will edit this post with the video when youtube is
done with maintenance.
http://www.youtube.com/watch?v=ozr1C6lL_DE
EDIT: With the synth provided above with ETN to get the syth to work with CN you would have to use 24.5g CN instead of AN. right? Everything else
should stay the same.
[Edited on 3-9-2010 by Murda]
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hiperion42
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 interesting
I tried to work out the numbers using kno3 as the nitrate
and a 23% HCL solution.
urea : technical 46-0-0 urea fertilizer.
kno3 : 13(13.6)-0-45
There is 136 g/kg kno3 which is 9.71 moles of nitrogen per kg kno3.
9.71 x 60.06 = 583.2 g urea /kg kno3
23% HCL @ d 1.18g/cm3 = 271.4 g/l
Molar mass HCL 36.46 g/mol
271.4/36.46 = 7.44 mol/l
9.71 moles = 1305.12 ml needed for 1kg kno3
for 1l 23% HCL there would be required 766g kno3 and 446.25g urea
I am not sure if just using more 23% HCL is going to do the trick without having to modify the addition by adding the nitrate and urea to the HCL.
So the yield for one liter 23% HCL is somewhat lower.
Shouldn't it be much lower or am i roughly correct?
.....ejuu....................................................................Ffg..............................g.............
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Murda
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From what I have heard kno3 has better yields than using AN with this method. I am unsure though with your concentration of HCL. I was wondering if
there any way to confirm UN rather than detonating it.
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hiperion42
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In any case more 23% HCL will be needed to roughly get in the same range of yields gotten with 31 % HCL.
I can't see any immediate substantial obstacles with the 23% route aside from needing more (HCL).
I will have to test this.
"I was wondering if there any way to confirm UN rather than detonating it."
I don't understand. Can you reformulate please.
[Edited on 23-3-2010 by hiperion42]
.....ejuu....................................................................Ffg..............................g.............
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Blasty
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Quote: Originally posted by hiperion42  |
"I was wondering if there any way to confirm UN rather than detonating it."
I don't understand. Can you reformulate please. |
I think that he means if there is a way to test that you have actually produced urea nitrate other than detonating it. In other words: is there a
chemical test for urea nitrate?
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ScienceSquirrel
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Urea nitrate melts at 152C
A simple melting point apparatus will suffice to characterise it.
Nice sharp melting point at 152 C to a clear liquid, you have pure urea nitrate!
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franklyn
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Metathetical double decomposition of selected salts wll obtain Urea Nitrate without
the need for liquid acids. I had made these observations
http://www.sciencemadness.org/talk/viewthread.php?tid=7553#p...
http://www.sciencemadness.org/talk/viewthread.php?tid=2642&a...
Urea Oxalate has trivial solubility but heated in a solution of Calcium Nitrate should
exchange ions precipitating Calcium Oxalate which is virtually insoluble leaving a
relatively pure solution of Urea Nitrate.
Mol weights , Urea - 60 , Oxalic acid - 90 , Calcium nitrate - 164
{(NH2)2CO}2 •H2C2O4 + Ca(NO3)2 •4H2O => CaC2O4 + 2 (NH2)2CO •HNO3
. . . . 120 . . . . . . . .90 . . . . . . . . 236 . . Formula weights
. . . . . 4 . . . . . . . . . 3 . . . . . . . . . . 8 . . . .parts by weight
We see Urea Oxalate is formed from 4 parts of Urea to 3 parts of Oxalic acid.
It can even be bought in that ratio for the equivalent of $ 2 a pound , $ 28 total for 14 lbs.
$ 14 for 8 lbs => www.soapgoods.com/Urea-Prilled-Commercial-p-751.html
$ 14 for 6 lbs => www.soapgoods.com/Oxalic-Acid-p-889.html
Ca(NO3)2 •4H2O is
$ 8 for 10 lbs => http://axner.com/calcium-nitrate.aspx
The ratio of 7 parts Urea Oxalate ( mol weight 210 ) to
8 parts Calcium Nitrate tetrahydrate ( mol weight 236 )
is practically stoichiometric. For 15 parts by weight of
materials one obtains 8 parts by weight of Urea Nitrate.
Equal to the weight of Calcium Nitrate tetrahydrate used.
.
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franklyn
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Unless you can buy it in 50 pound bags locally , for smaller quantities this is
as economic as you will find on line , 25 pounds for $ 30 delivered.
http://cgi.ebay.com/ws/eBayISAPI.dll?ViewItem&item=35026...
.
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Murda
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Not to bring up an old thread. But I read that UN burns vigorously. Mine does not burn but melts to a yellow color. I am thinking it's not UN?
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grndpndr
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Calcium nitrate or dbl salt?
@ Franklyn,Ive made inquiry to the folks(onebay) selling the 'calcium nitrate'
about its being Calcium Nitrate or a double salt Calcium Nitrate/Ammonium nitrate.It is in fact the latter double salt,I was sent a sample and label
from the original packaging.EBay Repackaging operation but economical and honest.The other link youve provided
'axner' stocks calcium nitrate terahydrate in its crystaline form rather than
prills and also quite inexpensive though I dont have the price w/shipping at hand.But depending on amounts between.60/80cents a lb less shipping.So
far I havent located Calcium Nitrate tetrahydrate as fertilizer
only the dbl salt w/AN.The only places Ive found so far are ceramic supply with the calcium nitrate terahydrate.Or does it make any difference to the
reaction?
[Edited on 16-5-2010 by grndpndr]
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franklyn
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Yes , especially on Ebay it certainly is prudent
to interrogate the seller as to composition.
I have not actually done this , but in solution
the chemistry must proceed as expected.
The only questions are the effects of impurties
so recrystallized before is indicated. Then there
is always the issue with temperature versus
concentration. I would just mix as much of
the more soluble salt Ca(NO3)2 as the water
will hold and adding in the required amount of
the Urea oxalate. That prepared in much the
same away by adding the oxalate to a cold
saturated urea solution. Urea sulfate would
also work.
.
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Murda
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| Quote: Originally posted by grndpndr | @ Franklyn,Ive made inquiry to the folks(onebay) selling the 'calcium nitrate'
about its being Calcium Nitrate or a double salt Calcium Nitrate/Ammonium nitrate.It is in fact the latter double salt,I was sent a sample and label
from the original packaging.EBay Repackaging operation but economical and honest.The other link youve provided
'axner' stocks calcium nitrate terahydrate in its crystaline form rather than
prills and also quite inexpensive though I dont have the price w/shipping at hand.But depending on amounts between.60/80cents a lb less shipping.So
far I havent located Calcium Nitrate tetrahydrate as fertilizer
only the dbl salt w/AN.The only places Ive found so far are ceramic supply with the calcium nitrate terahydrate.Or does it make any difference to the
reaction?
[Edited on 16-5-2010 by grndpndr] |
The yara fertilizer that I bought was not sitting out. She didn't have it
but she made a special order for me and had it shipped to her store. In the end I payed around 30 bucks for a 50 lb sack. 15.5-0-0. The only thing she
can't get for me is of course AN. all the agri stores around here have urea on hand readily. During the winter months you can also get urea a ice melt
 That was the good thing about working at a salt factory. First time I seen about
20 tons of urea.
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quicksilver
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Yara is a very high quality product but they do one thing that is infuriating. They put Mica in the product. This makes it difficult to work on large
batches with gravity filtration due to the plunging of the filter medium with the micron size Mica. One a small scale, this is not an issue what so
ever.
But if one were to want to pull all the various nitrates on a multi-kilogram scale there would be a need for a very large Buchner Funnel & loose
spun glass. Even then the problem would not be completely eliminated. Because in addition to plugging the filter, Mica doesn't "settle down" in the
solution easily. The slightest swirl and the crap is moving about.
You COULD attempt to pull the Mica & clay off 1st. By crushing prills, using 60C hot water, the sludge would settle down within the clay. If at
that point you were careful enough (or get a very well tuned vacuum hose), you could decant the clear solution and begin work without that headache. A
common 5gal bucket should hold enough Kg in solution to make it worth the effort. To work though a 50lb bag may take about a week but you'd get a
promising level of ammonium nitrate. However with the HOT water, you're pulling all he nitrates (nothing wrong with that, you paid for them).
Evaporation would yield an enormous mix, but pure on both counts. Using cold water isn't that feasible as it wont drop all the clay down to "hold" the
Mica in sediment.
I may just try it and write about what I could get from a specific unit weight, in terms of time vs. yield.
[Edited on 18-5-2010 by quicksilver]
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