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Author: Subject: Sodium nitrate/urea complex
godchem
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[*] posted on 24-4-2009 at 23:34


Would this work for preparation/synthesis of lead azide?

http://www.infernolabs.co.uk/filehos...c5f1a58739631f
Mirror: http://www.infernolabs.co.uk/filehos..._hydrazine.gif
(The link doesn't works for me)
---------
Fusing nitrate and urea at 150°-200° results in an azide salt? Hmmm....

BTW, you have to copy/paste the links to see the files.
--------
Yes!?
In a little chemistry book "Gefährliche Reaktionen" (English="Dangerous reactions") it said, that urea can react with nitrates to produce toxic azides.

I think sodium azide will be only produced as a byproduct.
But it can purified by making HN3, which can destilled at about 50°C.
I have not tried it yet. They said, that it will react with nitrate. Maybe it will work only with ammonium nitrate!? Or another nitrate?!

It the book it also said, that the power of AP can increased by mixing it with potassium permangante (losing of stability). And some other interesting things...

Maybe I can get another look in the book (it's not my book, I found the on my "intership" at a lab (I hope it's the right word) at a lunch break; they have some interesting chemicals in the lab, heheheeeee: nicotine, benzoyl peroxide,...)
----------

I read this on "Explosives and weapons forum"

http://www.roguesci.org/theforum/showthread.php?t=190&pa...




\"When I was young I played with legos, but now I am older and I play with atoms. - Alexander Shulgin\"

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[*] posted on 25-4-2009 at 12:27


Heating this mixture doesn't look like a good idea, since the urea sodium nitrate complex already explodes when heated. The complex is described by Werther in J. pr. 35 [1895], 60. The paper also describes the calcium complex, which first melts on heating, then give ammonicial, then acidic fumes and then on quick heating explode violently leaving behind CaCO3. Mg salt said to dec. similar. Attempts to form complexes with some others like KNO3, Ba(NO3)2, Sr(NO3)2 but these only crystallized from solution unchanged. But, heating a strong oxidizer together with urea at those temperatures to begin with. Ammonium nitrate and urea mixtures in preparation of guanidine nitrate have been known to explode also. Distilling HN3 is extremely dangerous and a set-up like that can easily get blown to pieces. Its Bp is also 37, not 47 deg. Undextrinated lead azide is also extremely hazardous since crystal breakage can detonate it under water.

Attachment: Werther.pdf (506kB)
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[*] posted on 15-6-2009 at 08:11


I tried it - it works :o
Basically I just mixed stochiometric amounts of NaNO3 and urea. Added just enough boiling water to dissolve all.
Upon standing in the fridge for 1 night I obtained a small amount of fluffy crystalline material.
Yields are crappy because of high solubility, so if anyone attempts to make this stuff, dont try to precipate it. Just let it evaporate.

The dry crystals burns violently and explode upon melting, very much like the famous yellow powder.
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[*] posted on 15-6-2009 at 08:34


The synthesis calls for the fusion of the dry reactants; you probably have a mixture containing urea nitrate. . .
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[*] posted on 15-6-2009 at 23:00


No Im talking about the original topic of this thread which is a complex/clathrate of urea with sodium nitrate. It is a well known fact that urea forms clathrates with many compounds and the NaNO3 clathrate is described by credible sources - go google it. And read Formatik's post btw. The clathrate is formed from very concentrated solutions.

God knows what you get by fusing urea with nitrate. Its a completely different subject. Maybe traces of azide are indeed formed. Being able to destill HN3 from it however is nothing but vague speculation. All we have to support that claim is an obscure amateur drawing.
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[*] posted on 16-6-2009 at 02:18


Sorry Taoiseach, I read one or two posts with NaN3 writ large. . .
The clathrate too is of interest and more likely to form than the azide.
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[*] posted on 7-12-2009 at 10:04
A long story


Zinc had started this thread
YeOldeImpurities provided this reference of it
http://www.sciencemadness.org/talk/viewthread.php?tid=7553#p...
- Following this lead I had posted this request in references , later edited.
http://www.sciencemadness.org/talk/viewthread.php?tid=10601&...
solo provided what it seemed to be
http://www.sciencemadness.org/talk/viewthread.php?tid=10601&...
gsd recognized the citation as this
http://www.sciencemadness.org/talk/viewthread.php?tid=10601&...
Sauron had relayed to me the following :
" I'm afraid that things are murkier than gsd thinks.
The way the Hauptwerke is organized and cited ( everywhere not just in PATR )
is by band and section then page.
If you tell me where in PATR is the citation, or tell me the exact citation
the way it is in PATR with funny brackets and all, then I can find it - I have the
Hauptwerke."
and Sauron provided the reference
http://www.sciencemadness.org/talk/viewthread.php?tid=10601&...
Attached below is the photocopied text - Gif 1. Translated below here _

Metal derivatives of Urea and of Urea with bases and salts
( also complexed with acids ):

- CH4ON2 + NaNO3 + H2O. Prisms. Dissolved in water.
Explodes when heated ( W., J. pr. [1] 35, 60 ). -
- 6 CH4ON2 + Ca(NO3)2. Crystal ( from alcohol ).
Melts under decomposition; explodes with rapid heating ( W., J. pr. [1] 35, 57 ).
-

Here I thought the citations were for the articles , turns out they are referencing
the citations , similar to Chemical Abstracts.

Formatik has attached this paper above ( W., J. pr. [1] 35, pages 57 & 60 inclusive )
Werther , J.Prakt.Chem ( Journal für Praktische Chemie ) (1) 35, 57-60
http://www.sciencemadness.org/talk/viewthread.php?tid=7553&a...

Taoiseach confirms the validity of the material having at least some
potential usefulness as a possible propellant or low grade blasting compound.

I just discovered that Urea Sodium Nitrate inclusion compound is regulated
and has industrial application. Who Knew ! _
http://law.justia.com/us/cfr/title21/21-3.0.1.1.7.2.1.14.htm...

__ __ __ __ __ __ __ __ __ __ __ __ __ __ __ __ __ __ __ __


Mind you I do not speak German, translation of the article text given here is my
interpretation of the literal translation by automated software ( Babelfish etc.)
- If there be errors please state any correction.
Quote :
A hot concentrated solution of equal mols of Sodium Nitrate and Urea , when cooled,
crystallizes as a double salt in long prisms. These are not efflorescent , begin to melt
at 35 º yet at 100 º are not entirely liquid. Decomposition occurs at 140 º, and with
stronger heating , exhibit the reaction of the complexes described thus far. The salt
can be heated in water without precipitating as the Silver complex does. Heated to
the melting point and then drowned , it does not decompose and is crystallized again.
If it loses water of hydration and is dissolved in water to re-crystallize , it dispropor-
tionates precipitating Sodium Nitrate first followed by Urea. If one dissolves this mix-
ture in a little hot water, it crystallize as the complex again.


. . . .NH2 • H2O
. . ./ . . . . . . . . . . . . . . . The explaination given in the text and the structural depiction
OC . . . . . . . . . . . . . . . in Federov on the left here shows that water of hydration
. . .\ . . . . . . . . . . . . . . . comprises the composition.
. . . .NH2 • NaNO3

Some facts on Urea Clathrate forms
http://www.chemistry.uoguelph.ca/soldatov/C7100/examsolution...
See pg 14 Urea Clathrates here _
http://homepage.univie.ac.at/jeanluc.mieusset/Supramolecular...
Encyclopedia of supramolecular chemistry Vol 1 - J. L. Atwood, Jonathan W. Steed
- Gif 2. ( excerpt attached - Channel inclusion compounds of Urea.gif )
Chemistry and Technology of Agrochemical Formulations - D. Alan Knowles
- Gif 3. ( excerpt attached - Urea inclusion compounds.gif )


I have thought of approaches to enhancing this Nitrate " clathrate " of urea, Urea forms
an adduct with H2O2 as well, so it seems possible that peroxide could substitute for the
H2O of the clathrate. Possibly instead of the single complex , a mixture of Urea•H2O2
and Urea•NaNO3 may be derived using H2O2 as the solvent.

- Axt on Urea Peroxide preparation
http://www.sciencemadness.org/talk/viewthread.php?tid=3214&a...
from the given reference: Method of preparation
- The crystals are prepared by the following procedures. A solution of Urea in
30% hydrogen peroxide, in the molecular ratio 2: 3, it is heated in a Pyrex dish for
a few minutes at a temperature of about 60 ºC. When cool it is transferred to a
crystallizing dish for slow evaporation and crystallization. Usually the crystals grow
in the shape of needles,
-
Cameo Chem Urea Peroxide MSDS
http://cameochemicals.noaa.gov/chemical/1682

CO(NH2)2•H2O2 , Carbamide Hydrogen Peroxide Adduct , CAS - 124-43-6
Beilstein Registry Number: 3680414 , Beilstein 3, IV, 94 ,
Solubility in Water: ~ 500 grams/ liter @ 20 °C , Solution pH: 5.0 - 6.5
At about 40°C solution disassociates becoming turbid as H2O2 decomposes.
100 grams of Methylol dissolves 21.8 grams Urea @ 19.5 ºC
I guess the peroxide adduct has similar solvation in MeOH.
Solvates with disassociation , in alcohol , acetone or ether.
Forms detonable shock sensitive solution with acetone, particularly with ether.
Anhydrous Decomposition Temperature: 60 ºC
Flash Point: 60 ºC

____________________________________________________


There is also an isomer of the Nitrate anion, called Peroxynitrite,
which is more liable to disruption , the shape is linear and more amenable to
cage formation , which may yield a useful ' green ' primary.
While I believe the following is accurately related it is a compilation from various sources.
The classic method of synthesis is :
Cold weakly acid solutions in the range of pH 4.0 - 9.0 ( optimally in the region of 6.5 to 7.0 )
activate NaNO2 in H2O2 to react, forming peroxynitrite ( ONOO- ) which is stable at high
basic pH ~13.1 At pH below 4.0 ( 3.0 - 3.5 ) nitrite is instead oxidized to nitrate by peroxide.
The anion is stable, but protonated to acid (pKa = 6.8) it rapidly decays to nitrate at 25 ºC
Peroxynitrous acid can persist for longer only below - 80 ºC.


The idea then is this, Urea is first to be crystallized as an adduct with Hydrogen Peroxide.
This then is to be mixed with Sodium Nitrite in solution producing Sodium Peroxynitrite
in situ, the formation of the hypothetical Urea•Sodium Peroxynitrite complex or even a
Urea•Hydrogen Peroxide•Sodium Peroxynitrite complex would there upon be crystallized
as the desired resulting product.

The following chart naively shows the permutations which may arise but are untested.
1. The known Clathrate. 2.Supposed H2O2 variant. 3. Supposed Peroxynitrite variant.
4. Clathrate of both H2O2 & Peroxynitrite

Complexes . . . 1 . . . 2 . . . 3 . . . 4

Urea . . . . . . . . . . . . . . . . . . . . . .
NaNO3. . . . . . . . . . . .
H2O . . . . . . . . . . . . . . . . . . .
H2O2 . . . . . . . . . . . . . . . . . . . . . .
Sodium
Perxynitrite . . . . . . . . . . . . . . . . . .


NaNO2 , Sodium Nitrite , CAS - 7632-00-0
Hygroscopic and very soluble in water, but little soluble in most organic solvents.
Solubility in Water: 850 grams / liter , @ 20 ºC , Solution pH: 9.
Solubility in DMSO is , 20 gms per 100 cc @ 25 ºC
!00 grams of Methylol dissolves 4.4 grams NaNO2 @ 19.5 ºC
( Methyl Nitrite requires much longer induction period than peroxynitrite formation )

" Peroxonitrous acid is an acid of medium strength (pKa = 6.8), which rapidly decomposes
in neutral conditions ( half life = 1 sec at pH = 7.4 ). Peroxonitrite can isomerize to nitrate
or decompose to nitrite and dioxygen. Isomerization to nitrate is a major pathway in acidic
media. The mechanism for the decomposition is still in doubt, but it is believed that HOONO
homolyzes to give the NO2/OH radical pair. Peroxonitrites are stable in basic solution.
Peroxonitrous acid is formed as a yellow intermediate in the reaction of nitrite solutions with
hydrogen peroxide that leads to nitrates. HNO2 + H2O2 ONOOH + H2O
It is a powerful oxidizing agent, stronger than nitric acid or hydrogen peroxide."

The following is excerpted from Encyclopedia of Inorganic Chemistry PDF page 2986
" After much confusion over the past 10 years, it is now clear that in aqueous buffer
in the absence of another reactant, peroxynitrous acid ( ONOOH, pKa = 6.8 ) undergoes
two subsequent reactions, isomerization to produce nitrate ( yield 71% ) and homolysis
to produce hydroxyl radical (•OH) and nitrogen dioxide (•NO2) ( yield 29% ). Both of
these species are highly oxidizing and much speculation has been directed toward their
involvement in biologically damaging actions of (•NO) production. However, under biological
conditions, peroxynitrite will undergo at least two other reactions long before isomerization
and homolysis, two-electron oxidation of cellular thiol ( forming the sulfenic acid and nitrite ),
and reaction with CO2 to form the nitrosoperoxocarbonate anion (ONOOCO2− ).
This extremely short-lived species undergoes either heterolytic scission, producing nitrate
and CO2 ( 67% yield ) or homolytic scission producing (•NO2) and the carbonate anion
radical (•CO3− ) ( 33% yield ).
"

Peroxonitrite Physical Properties Data
http://www.caymanchem.com/pdfs/81565.pdf
Peroxonitrite MSDS
http://www.caymanchem.com/msdss/81565m.pdf
Other references
http://pubchem.ncbi.nlm.nih.gov/summary/summary.cgi?cid=1048...
Investigation of the interaction of sodium nitrite with hydrogen peroxide in aqueous solutions by the chemiluminescence method
http://www.springerlink.com/content/7472nk4rj4g06665
http://resources.metapress.com/pdf-preview.axd?code=7472nk4r...

solo has very kindly provided this important reference paper
The chemistry of peroxonitrites
J. O. Edwards; R. C. Plumb
Progress in Inorganic Chemistry 41:599–635; 1994
http://ifile.it/5typia8

solo , also very kindly provided this other reference paper
Determination of optimal conditions for synthesis of peroxynitrite by mixing acidified hydrogen peroxide with nitrite
Saha A, Goldstein S, Cabelli D, Czapski G. - Free Radic Biol Med. 1998 Mar 1;24(4):653-9
Department of Physical Chemistry, The Hebrew University of Jerusalem, Israel.
Abstract, in part:
High yields of peroxynitrite were obtained at room temperature using an efficient double
mixer where acidified peroxide was mixed with nitrite; after an appropriate delay, the
reaction was quenched with strong alkali. An excess of more than 10% of H2O2 over nitrite,
or vice versa, is sufficient to get ca. 85-90% of peroxynitrite, almost free from nitrite or
H2O2, respectively. The results also suggest that conventional use of ice-cold solutions
of the reactants and the alkali solutions is not required if an efficient mixer and appropriate
quenching times are available.


.

Hauptwerk.gif - 11kBChannel inclusion compounds of Urea.gif - 37kBUrea inclusion compounds.gif - 30kB
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[*] posted on 7-12-2009 at 11:43


I mixed 20 g of urea and 28 g of NaNO3 (0. 33 moles of each) in around 40-50 ml of water (mixed the powders, and while heating on a hot plate adding water, when the solution boiled I removed it from the heat). Now I placed the solution in a freezer to cool it. I will report the results when it precipitates andd I filter it.


Edit:
After cooling the solution to -5 C only a small amount of powder preciptated. It seems that I made the solution too dilute, so now I placed it on a radiator to slowly vaporate off the excess water.

Also I scooped with a metal spoon a small amount of the precipitated powder and while still weat heated it with a gas burner. The powder first dissolved, dried and decomposed rather unenergeticaly. I will test it again when the solution dries.

[Edited on 7-12-2009 by Zinc]




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[*] posted on 9-12-2009 at 04:05


Excellent work franklyn !

Replacing the hydrate water with H2O2 sure looks very promising but might be difficult to achive practically. The urea*H2O2 adduct is stable in solution at low temperatures only, whereas the clathrate formation needs hot saturated solutions so that's getting us nowhere.

Sodium, magnesium and calcium nitrate are insoluble in ethanol. Mr. Werther seems to suggest that the complexes are soluble however - this is suprising but quite possible, considering the cage-like structure of the complex and the solubility of urea in ethanol. So it *could* be as simple as: 1.) Make the nitrate-urea-clathrate 2.) dissolve in minimum amount of warm ethanol 3.) add calculated amount of H2O2 4.) put in freezer and pray to receive the hoped-for urea peroxide-sodium nitrate clathrate :)
Other than that, one might wet the dry compound with the calculated amount of H2O2 and put in a dessicator over H2SO4. It could take several month to dry tough (if it drys at all!)

Also we need to find a patent describing a high-yielding synthesis of the urea-sodium nitrate complex, now that we know it has comercial applications. Mixing hot saturated solutions give only pathetic amounts, and slow evaporation causes NaNO3 and urea to crystallize seperately. Using hot saturated solutions of urea in ethanol and NaNO3 in water could help improve yields. Slow evaporation does work with the Ca(NO3)2 adduct tough, according to Werther. This stuff might be even more interesting as both ingredients are available as fertilizers.

>The crystals are prepared by the following procedures. A solution of Urea in
>30% hydrogen peroxide, in the molecular ratio 2: 3, it is heated in a Pyrex dish for
>a few minutes at a temperature of about 60 ºC. When cool it is transferred to a
>crystallizing dish for slow evaporation and crystallization. Usually the crystals grow
>in the shape of needles, -

This doesn't look right. I've prepared the adduct by fist making a saturated urea solution (40°C seems fine) and cooling this down to 0°C. The solution is then filtered into the stochiometric amount of cold H2O2 (-5°C) and the mixture put in the freezer for an hour or so. The precipate is washed with alcohol and a few drops of citric acid solution to stabilize the adduct. It is stable below pH 4 only. It never came out in the form of needles but rather coarse chunky transparent crystals or plates. If you get white needles then thats urea precipating either due to over-saturating the solution or decomposing the complex. At 60°C you might get mostly urea and little peroxide as this adduct is thermally unstable. The dry stuff gives off white fumes at about 60°C. A lit match put into the fumes give quite some firework. The reaction is so violent that it cannot be accounted for by O2 only. Rather the fumes seem to contain a lot of radicals.
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[*] posted on 10-12-2009 at 10:02


The formula for " Greek Fire " is a matter for speculation as it is unknown today.
The supposition that Urea•NaNO3 composition is or could have been the secret
ingredient really begs the question. What peaked my interest initially is that NaNO3
has seen use for a long time , a thousand years at least. That the pyrotechnic
nature of its combination with Urea has remained obscure is most curious.
__ __ __ __ __ __ __ __ __ __ __ __ __ __ __ __ __ __ __ __ __ __ __

There are many vicissitudes in just how components are brought together in a
solution crystallizing to form a complex. Myriad subtleties and nuances require
a working procedure to facilitate formation. How component ingredients are
brought together, in what circumstances and ambient conditions.
Quote:
Excerpted from Crystallization 4th ed. pgs 142-143
The solute and solvent of a solution may, and frequently do, combine to form
one or more different hydrates. For instance Na2S04•10H2O melts at 32.4 ºC
to give a saturated solution of sodium sulfate containing a suspension of the
anhydrous salt.

- Even their description in references can be ambiguous and misleading -
Quote:
Excerpted from Crystallization 4th ed. J.W. Mullin , pgs 90-91
Many methods of solubility expression can lead to the use of the potentially
misleading term ' percentage concentration '. For instance, an expression
such as ' a 1O per cent aqueous solution of sodium sulfate ' without further
definition, could be taken to mean anyone of the following:

10 g of Na2S04 in 100 g of water
10 g of Na2S04 in 100 g of solution
10 g of Na2S04•10H2O in 100 g of water
10 g of Na2 S04•1OH20 in 100 g of solution

If 10 g of anhydrous Na2S04 in 100 g of water were the intended description
of the solution concentration, this would then be equivalent to :

9.1 g of Na2S04 in 100 g of solution
20.6 g of Na2S04•10H2O in 100 g of solution
26.0 g of Na2S04•10H2O in 100 g of water

which gives some measure of the magnitude of the possible misinterpretation.
To make matters even worse, the term ' percentage concentration ' is often
applied on a volume basis, e.g. 10 g of Na2S04 in 100 mL of water, of solution,
and so on.



Taoiseach you say " Urea•H2O2 adduct is stable in solution at low temperatures only "
Yes once formed, the Urea•H2O2 adduct nevertheless forms from cooling a hot 60 ºC
saturated solution of Urea in 30 % H2O2, the same as the Urea•NaNO3•H2O clathrate.
See the cited paper http://www.sciencemadness.org/talk/files.php?pid=59337&a...
From it's description Urea•NaNO3 also disproportionates in solution after it's lost water
of hydration. The resulting solution of Urea to NaNO3 in a 1:1 molar ratio needs to be
again saturated ( hot ) to reform and re-crystallize as the clathrate. This proposed
complex may form simply from a hot solution of Urea, NaNO3, in 30 % H2O2. If one
were to attempt to blend NaNO3 with the already formed Urea Peroxide adduct, then
DMSO or perhaps 2 :1 molar ration of Urea to Choline Chloride ( deep eutectic solvent )
may serve as anhydrous solvents. This anhydrous approach may better serve the
investigation of the formation of Peroxynitrite in situ, Urea Peroxide + Sodium Nitrite.
For mixed solvents MeOH is preferable to EtOH as it has higher solvation of Urea ,
100 grams of Methylol dissolves 21.8 grams Urea @ 19.5 ºC. ( for NaNO3 it's very poor )
It also tends toward lower pH ( world's worst acid ). As I say, there are many vicissitudes.


P.S.
- How's it coming along Zinc , we are all ears.


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[*] posted on 10-12-2009 at 13:21


Well after most of the water evaporated, I took a small sample of the mix (almost dry) and heated it with a gas flame. The solid first melted and started decomposing. Only a few small pops were heard. The second time I did that not even the small pops were heard. So it seems that the complex cant be isolated by slowly evaporating off the water. I will add some hot water to dissolve the mix and cool it. Perhaps that will work.



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[*] posted on 18-12-2009 at 07:07


To provide a baseline for comparison of subsequent ingredient variations
I prepared a 1 : 1 molar ratio of agricultural prill Urea and Nitrate of Soda
as follows:
These I have not assayed, the Urea appears dirty although the product
specifications indicate nothing is admixed.
http://agr.wa.gov/PESTFERT/Fertilizers/FertDB/prodinfo.aspx?...
The NaNO3 is similarly suspect since the fertilizer code 15-0-0 indicates
15 % nitrogen when it should be 16-0-0 , 16 % ( 14 / 85 , mol weight of
Nitrogen divided by mol weight of NaNO3 ) indicating hydration or inactive
content. Product specifications indicate nothing is admixed
http://www.cdfa.ca.gov/egov/is/fert/fert2.asp?ID=5776
http://www.bonideproducts.com/lbonide/msds/msds915.pdf
* Note : 15.5-0-0 is Ca(NO3)2 , 15-0-14 is Nitrate of Soda & Potash

I weighed 1/5 mol of each,
12 gm Urea + 1.5 gm extra , and 17 gm NaNO3 + 2.5 gm extra, the slight
excess to compensate for what is lost in filtering. I dissolved each by adding
the minimum of boiling water, hot filtered separately and let stand. Despite
anticipating loss in processing I guess the amount soaked by the paper is
at least 5 gm of each ( without weighing the dry paper with filtrate which I
did not do ). This has had the effect of skewing the molar ratio from 1 : 1
to 2 : 3. Urea left some slight insoluble residue and the solution retained
a tea like tinge , the NaNO3 left none. After first cooling to room temperature
and even after keeping in the refrigerator there was no evidence of either
crystallizing. I poured both solutions together and set it on the steam radiator
to dry overnight. The solute collected in the shallow plastic pan as a light tan
colored solidified sheet, weighing 24.8 gm.
Subjected to the flame of the stove, a small sample merely melted as a
syrupy bead and the gas flame took on the characteristic orange color
indicating the presence of Sodium.
I partially dissolved the product again with boiling water and heated this
further until it all dissolved and set to cool on a plastic pan. No crystallization
occurred and I placed the liquor to dry on the radiator. After a few hours
the congealed mass reappeared and I again applied the stove flame test
with identical result.

The unimpressive result can be attributed to the hydration in this clathrate.
The enthalpy of vaporization of water is very high, and whatever heat is
generated from the combustion of Urea and Nitrate is absorbed by the water
of hydratation turning to steam. A dry mixture of Urea and NaNO3 alone is more
energetic , a clathrate just provides an ideal mixture at molecular scale. The
reason I thought to substitute the water with Peroxide.


Mol wt. of H2O2 is 34 gm , Mol wt of water is 18 gm , in a ratio of , 2 : 9 ,
( 2 X 34 = 68 ) and ( 9 X 18 = 162 ) , 68 + 162 is 230 , then 68 /230 is ~ 30 %
~ 100 volume H2O2 solution , 230 gm of which contains 2 mols H2O2.

Volume measure means 1 volume of liquid yields 100 volumes of oxygen gas
measured at 0 ºC in atmospheric pressure . 30 % concentration is 300 gm
of H2O2 in a 1000 gm of solution. 300 divided by 34 ( mol wt of H2O2 ) is
8.8 mols H2O2 in a kilogram of solution. 8.8 mols multiplied by 230 gm of solution
( 8.8 X O.23 ) is 2 mols of H2O2 content.
2 mols of H2O2 yields 1 mol O2 ( 2 H2O2 -> 2 H2O + O2 ) and so
8.8 mols H2O2 yields 4.4 mols oxygen. A molar volume of gas at 0 °C and
atmospheric pressure occupies 22.4 liters , then 4.4 X 22.4 = 98.6 volumes
( note the discrepancy from 100 is due to rounding the mol ratio of H2O2
and H2O to integer amounts , 2 : 9 )


Mol wt. of Urea is 60 gm , Mol wt. of NaNO3 is 85 gm , 2 mols of each is
( 2 X 60 = 120 gm ) and ( 2 X 85 = 170 gm )

1/10 of each ( 0.2 mol ) is 12 gm Urea , 17 gm NaNO3 , 23 gm 30 % H2O2 solution.

solution of Urea is 250 gm per 100 gm water at 60 ºC
250 . - . 12 gm Urea
100 . . . 4.8 gm H2O

solution of NaNO3 is 124 gm per 100 gm water at 60 ºC
124 . - . 17 gm NaNO3
100 . . . 13.7 gm H2O


18.5 gm total water required , 23 gm of 30 % H2O2 contains 16.2 gm water
and solvation provided by the 6.8 gm H2O2 is practically the same as water.

Mix in and dissolve the solutes into the 30 % H2O2 solution , a total 52 gm.
Maintain at 60 ºC to evaporate until volume has decreased and crystallization
has begun. You can then add 20 gm or so of an alcohol to speed this up.

This is what I am doing , and will report on the result.

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Taoiseach
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[*] posted on 18-12-2009 at 16:21


Interesting find. The paper mentions that the complex looses water easily and then breaks down into an intimate mixture of urea and NaNO3. I wonder if the explosive properties pertain to the complex at all, or if it is just the transient formation of the complex which then brings about a mixture of the *dry* compounds at molecular level. When I did the experiment I did get a crystalline precipate from the cooling solution but the yield was terribly small. I filtered the few small crumbs and dried with a heat gun. Maybe that removed the water of hydration. The resulting dry stuff was definetly energetic, altough it was too little to do any more specifc tests than igniting it.

Anways keep us posted as to the results of your peroxide experiment!
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[*] posted on 24-12-2009 at 15:41


I'll cut to the chase and say the Urea H2O2 NaNO3 composition exhibits no
pyrotechnic traits , at least for me.


What a pain it has been to produce some Urea reasonably free of contaminants.
Without detailing the travails I recommend the bought material is first dry sifted
in a pan to separate prill beads from a darker granular content , sugarlike like in
appearance , which is thrown away. This has much clay in it which is admixed to
provide nucleation for prills to form as it is sprayed in the cooling tower when it
is made. The starting water solution should be lukewarm , filtered like that ,
( Filtering was done through coffee filters on a wire mesh sink drain screen which
serves me as well as any Buchner funnel ) then placed into the refrigerator freezer
for perhaps 20 minutes at most and retrieved before it can freeze. Tip the container
to allow the free running colored liquid to drain to reveal a conglomerate of needles
which can pass for white. These are harvested and the drained liquid is returned to
the needle mass which remained colored, warmed again to dissolve and placed in
the freezer to repeat the process.
Stay away from hot water solutions as this instead of a discretely crystallized
material turns into slush as it cools. Alcohol contrary to expectation provided no
advantage in clean crystallization. http://pubs.acs.org/doi/abs/10.1021/je60054a020
http://ntrs.nasa.gov/archive/nasa/casi.ntrs.nasa.gov/1989001...

The stipulated 15-0-0 Nitrate of Soda requires some adjustment to the weighed
molar proportion since that is on the basis of 16-0-0 which indicates pure content.
So ( 15 / 16 ) = 0.9375 , therefore the necessary 17 grams of NanO3 divided
by 0.9375 indicates 18.13 grams of the 15-0-0 material are to be used.

- Proceedure

Solutions were warmed in a microwave oven for no more than 15 seconds at any
one time. Temperature was determined by an electronic infrared thermometer.
Weights were assessed on an electronic scale which reads to 1/10 th of a gram.

23 grams H2O2 at ambient temperature is added to the 12 grams Urea and stirred
untill dissolved. 18.1 grams 15-0-0 NaNO3 is stirred into this , becoming dissolved
only when a slight excess of H2O2 is added to the already calculated amount and
the solution is warmed to ~ 55 ºC .

The solution is at first cloudy then clears without evolution of gas.
Some small insoluble precipitate collects on the bttom which I scoop out and
test on the stove burner. The material passively gives an orange color to the
flame indicating Sodium is present. This could be inert content in the 15-0-0
NaNO3 , possibly Na2CO3. The H2O2 used also contains some very minor amount
of phosphoric acid. The solution is left to dry in a styrofoam tub on a radiator
which maintains the solution at 55 ºC. After 3 hours submerged crystals appear ,
and I witness a barely noticable effervescence of occasional pin point bubbles
breaking on the surface. I determine this is caused by being heated when I
move the tub away from the heat and return it again to recieve heat.

I then place the tub inside the freezer. After 20 minutes the contents have
solidified but very soon begin to show a small amount of runoff. I pour out a
few drops of this liquid into a saucer and add to this a few drops of Sodium
Hypochlorite solution and considerable effervescence results indicating this
is H2O2 of at least original strength ( I cannot tell if it has been concentrated ).
After a few moments I pour out more runoff that has collected and break up
the mass of needles to air dry better. I add more Sodium Hypochorite to the
liquid I poured out , producing much effervescence and it took on a barely
discernable pinkish tinge.

I set the tub with the stirred needles back on the radiator. After some minutes
I check the progress and discover that the material has melted exuding fluid ,
I again place the tub back into the freezer to harden the contents. Some time
later I take it out and place it in sunlight on the window sill where it remains
for the afternoon. The material remains solid and I harvest the content to
weigh , yielding 27.6 grams. This is visually indistinguishable from urea itself.
Applying a small sample to the stove burner shows the characteric orange
flame except now there is some hint of flaring. I apply a small sample to a
saucer and nearby a small puddle of Sodium Hypochlorite solution then tilt
the saucer to run both together. Immediately there is strong effervescence
indicating that considerable H2O2 is present in this material , and the sample
entirely dissolves. To cover every possibility, I place a small amount inside
of folded Aluminum foil on top of a heavy steel plate covered by a light steel
plate which I hit with hammer blows three seperate times without effect.

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[*] posted on 27-12-2009 at 10:18


A History of Greek Fire and Gunpowder - Partington, J.R.
Cambridge, England: W. Heffer and Sons, LTD, 1960.

A comprehensive study on the origins and use of incendiaries and gunpowder weapons
from the earliest times. The first chapter looks at incendiaries, and includes investigation
into the ingredients of Greek Fire. The final chapter is on saltpeter. English chemist and
historian J.R. Partington traces their origins to Assyrian bas-reliefs from the 9th century
B.C., and even finds hints of them in the Old Testament ( Proverbs 26:18 ), the advent
of so-called Greek fire used by the Byzantine fleet to defend Constantinople against Arab
attackers in the 7th century, and later against the Crusaders. One of history's first secret
weapons, Greek fire is poorly understood today, contemporary accounts describe nozzles
spouting a fiery liquid that would burn even on the surface of the sea. Experts have tried
to determine the exact nature of the substance without reaching any definitive conclusions
( the recipe has been lost in time ). Partington offers his own theories about one of the
great mysteries of premodern warfare, he thought that Greek fire was made of a distilled
petroleum fraction and other ingredients but not saltpeter. He also traces the advent of
gunpowder to 11th-century China , exploring the legend of supposed inventor Black Berthold
( a mythic figure, says Parrington ) and examining the development of firearms in Europe,
the Middle East, and China.

James Riddick Partington was a distinguished historian of science and a
professor of chemistry at the University of London.
http://www.worldcat.org/isbn/0801859549
http://www.amazon.com/gp/product/0801859549

An account of " Greek Fire " is in COPAE by Tenney pages 32 - 34
Available in the SciMad Library _
http://library.sciencemadness.org/library/books/the_chemistr...

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