KETO-RDX and DIKETO-HMX

[VasiaPupkin]

I have found an interesting method of keto-rdx preparing in Journal of organic chemistry 2000-2002.
This method is based on dinitrourea.

1. N,N'dinitrourea hexamethyleneteramine nitrate salt (DHN): 21g of concentrated nitric acid mixed with 21g 20% oleum. Mixture cooled at -5-0C
and add 6g of urea by small portion with stirring. After 30 min of keeping at -5-0C this mixture was poured into hexamine solution (14g hexamine in 50ml water) at 15-20C.
Precipitate was washed with ethanole and diethyl ether and dried at room temperature. Yeild 19.8g (56%) m.p. 98C.

Dinitrourea also might be prepared without oleum:
To 40ml mixture of 50% nitric and 50%sulfuric acid (d=1.756) add 10g urea by small portions at 0-5C. After 25min of keeping of that temperature mixture was cooled to -12C. Precipitate was filtered.
Crude product (24g) has 80% of dinitrourea.

2. To 24ml mixture of 50% nitric and 50%sulfuric acid (d=1.756) add 14.12g of DHN at -5-0C with stirring. Stir mixture 30min at 20C. Then mixture was poured into 40g ice with water.
Precipitate was filtered, washed with ethanole and diethyl ether and dried at room temperature. Yeild 8.4g (89%) of 1,3,5-trinitro-1,3,5-triazacyclohexane-2-one m.p. 170-180C. (200-201C with dec. from nitromethane ).

What do you think if I'll take an dinitrobiuret (O2NNH-CO-NH-CO-NHNO2) instead of DNU?
Is it possible to make a diketo-HMX (2,4-dioxo-1,3,5,7-tetranitro-1,3,5,7-tetracyclooctane) by this method?

[fulminans]

What about the condensation of DNU with 1,1,3,3-tetraethoxypropane? You would get a rather stable NU-Derivative with excellent performance according to the literature.

Bibliographic Information

Synthesis and properties of tetranitropropanediurea. Hong, Guanlin; Zhu, Chunhua. Xian Mod. Chem. Res. Inst., Xian, Peop. Rep. China. Proceedings of the International Pyrotechnics Seminar (1991), 17th(Vol. 1), 193-5. CODEN: PPYSD7 ISSN: 0270-1898. Journal written in English. CAN 116:197173 AN 1992:197173 CAPLUS

Abstract

Tetranitropropanediurea (I) was prepd. by nitration of propanediurea with Ac2O-HNO3. The properties of I are reported (e.g., d. 1.98 g/cm3, activation energy of thermal decompn. 263.6 kJ/mol, detonation velocity 9034 m/s). I was prepd. from 1,1,3,3-tetraethoxypropane with urea in acidic medium.

Indexing -- Section 50-2 (Propellants and Explosives)
Section cross-reference(s): 21

Nitration
(of propanediurea)

Explosives
(tetranitropropanediurea, prepn. and properties of)

42563-68-8P
Role: RCT (Reactant); SPN (Synthetic preparation); PREP (Preparation); RACT (Reactant or reagent)
(prepn. and nitration of)

115029-36-2P
_
Role: PRP (Properties); SPN (Synthetic preparation); PREP (Preparation)
(prepn. and properties of)

Supplementary Terms

nitropropanediurea prepn property; propanediurea nitration

[nuclearattack]

What is the difference, from the performance point of view, between keto-rdx and the classic RDX?
It seems to be an easy way to make a powerfull HE but it can be interesting only if the performance is similar to RDX.

[Sylla]

With plus one oxygen and minus two hydrogen keto-RDX has a better oxygen balance. I'm pretty sure it will release more energy too and have and higher VOD...

I'll do some theoretical expectations later at home. I must say it seems to be a fucking good HE :]==]

[EDIT - Warning: APPROXIMATIONS] Results are close to those of RDX (for same density), I've approximated the VOD (at d=1.82) to be around 8618m/s. Notice that if TRIKETORDX existed it would have been (only) about 7921m/s...

[VasiaPupkin]

The measured perfomance of keto-rdx in 4% higher than HMX. Also you can read something in US pat. 5,391,736

[fulminans]

The measured perfomance of keto-rdx in 4% higher than HMX. Also you can read something in US pat. 5,391,736

New Data for K6 confirm it's of similar performance like HMX, but not superior. One advantage of K6 is the ease of preparation, no need for Acetanhydride!!!. One major drawback is the impact sensitivity. Can anybody get more Literature, I just have the appropiate links and abstracts so far.

Relationship between crystal shape and explosive properties of K-6. Ritter, H.; Braun, S.; Schafer, M.; Aerni, H. R.; Bircher, H. R.; Berger, B.; Mathieu, J.; Gupta, A. Institut Franco-Allemand de Recherches de Saint-Louis (ISL), Saint Louis, Fr. International Annual Conference of ICT (2001), 32nd(Energetic Materials), 91/1-91/14. CODEN: IACIEQ ISSN: 0722-4087. Journal written in English. CAN 135:333014 AN 2001:605107 CAPLUS (Copyright 2004 ACS on SciFinder (R))

Abstract

2-Oxo-1,3,5-trinitro-1,3,5-triazacyclohexane (K-6) was synthesized according to the two step pathway over 2-oxo-5-tert-butyl-1,3,5-triazacyclohexane. Checking different nitrating agents, the best yields (61%) have been achieved with nitronium tetrafluoroborate (NO2BF4) in acetonitrile. Different recrystn. procedures from Et acetate resulted in two types of crystals K-6-D and K-6-I, with differering behavior. Both products show very similar IR, MS and NMR spectra; by means of NMR spectroscopy the hydrolysis reaction of K-6 in d6-DMSO and subsequent disproportion of the intermediate formed was shown. Thermoanal. methods (DTA/TG, DSC and Hot Stage Microscopy HSM) showed a two-step decompn. at about 200° for both types and some reactions of the K-6-D crystals due to crystal phase changes or tension effects. The NOx development of K-6 at 90° detected by chemiluminescence (CL) did not exceed the range of known explosives. Some differences in the behavior during handling safety testing were also found. BAM friction sensitivity testing of K-6-I results in a reaction threshold of 48 N load, whereas in two trials the reaction was violent enough to break the porcelain plate. K-6-D already reacts at 36 N, but its reaction is only moderate. In the BAM impact testing all K-6 samples except one coarse crystal fraction reveals a pronounced sensitivity of 2.0 Nm, which approaches primary explosive behavior. Only the large crystals from DPA (4.5 Nm) are in the sensitivity area of HMX. The explanation of this fact can be found in different crystal shapes. Whereas X-ray diffraction investigations showed an orthorhombic elementary cell of the same size for both K-6 types, clear differences in the shape of the crystals could be shown with SEM. Whereas K-6-D consists of well built flat crystals, K-6-I has the structure of agglomerated thin plates. The explosive properties of K-6 have been assessed by comparing its measured detonation velocity and brisance with other explosives.
K-6 exceeds obviously the performance area of TNT or PETN, but a level higher than HMX cannot be achieved.

[mongo blongo]

How is this DHN salt going to form when the DNU will hydrolize in the water present in the Hexamine solution? This sounds very interesting by the way.

[VasiaPupkin]

First of all DNU needs some time (a few hours) to decompose to nitramide in water solution.
The reaction of DHN salt formation is a quick therefore DNU dont have a time to decompose.
It looks like DNU is sufficiently stable in form DHN salt.

[mongo blongo]

Ah right. I was under the impression that it happens quickly. One more thing, is it necessary to have the acids present when forming the DHN salt (as they are with the oleum method) or does the Hexamine need to be hydrolized by the acids? I was just thinking if it is needed to bother isolating the DNU when using H2SO4 instead of oleum.
I can't see any reason why K-RDX can't be done using just nitrate/H2SO4.
No HNO3 needed ;)

[VasiaPupkin]

I think acids is have to be when forming DHN-salt. Because DHN is a complex salt with nitric acid (CH2)6N4.OC(NHNO2)2.HNO3. And I think acid is decrease a DHN solubility in water.
Yes I also propose to use AN/H2SO4 mix for nitration in practice.
But its absolutely not necessary to separate DHN salt. You can make it in "one pot". Its even more convenient method for acid economy espesially if you use no oleum.
Yeild is also good but product might to be with more RDX impurity.
I have made some calculations and it must looks like:
10 g urea dissolve in 130ml H2SO4/HNO3 50:50 (or calculated amount AN/H2SO4) with small portions and with stirring at -5C.
After 30min -5C keeping add slowly 23g of Hexamine and slowly treat it to +15C. After 30min pour it into ice and collect precipitate.
Something like that.

[Rosco Bodine]

VasiaPupkin ,

That JOC article is very interesting . Perhaps someone could upload
the article to the FTP .

I have urea nitrate and hexamine dinitrate on hand . Do you think
that these acid salts could be used to any advantage , perhaps
to improve the yield of the DHN precursor , or possibly to
entirely eliminate the first step of isolating the DHN , producing
a good yield of the end product directly in one step ?

I was thinking that using the acid salts eliminates the extra heat of
formation that would otherwise be problematic when urea or hexamine
contact the acid . Converting the urea nitrate to nitrourea first may
also be a possibility , and that would eliminate a molecule of water
from the process at the beginning . But I do not know whether or not
the nitrourea would decompose upon contact with the acid mixture ,
as opposed to the desired reaction of forming nitrourea sulfate and
then nitrating further to dinitrourea .

The 50/50 acid mixture at the density described would consist of highly
concentrated nearly anhydrous acids . So any strategy for reducing the
byproduct water produced by any of the reactions involved should be helpful .

[mongo blongo]

I'm having a hell of a time forming DNU! I'm using NH4NO3/H2SO4 and pure urea. The H2SO4 is drain cleaner (91%) which was boiled untill very dense fumes evolved. The NH4NO3 and urea were so dry that you could die of thirst if you looked at it for too long.
Acids were cooled to -5°C and kept there during the urea addition and for 30 mins after. The reaction by this time was cloudy with suspended bubbles and what looked like small white particles.
It was cooled to -15°C and looked the same as before cooling with no crystal mass.
I tried reacting it with the Hexamine but I just got the faint smell of formaldehyde and a bubbling (not formaldehyde) when the mix was swirled. Again no precipitate formed even when cooled to 0°C.
I'm going to try using more H2SO4 next time because I can see small bubbles suspeneded in the nitration reaction indicating a possible decomposition product of the DNU (CO2). I might even try the reaction at -10°C.
Anybody tried this with a nitrate before?

[VasiaPupkin]

I think urea nitrate and hexamine dinitrate could be useful only if you use AN/H2SO4 mix.
Its allow to use low AN amount and decrease NH4+ level in nitration system to make experiment more close to method from JOC.
The heat rising is not so significant due to large acid amount.
AFAIK nitrourea sulfate is not exist (even in nitration mixture).
The only way to rise a yeild KETO is to use nitrourea (synthesed beforehand) instead of urea.

Hmmm. No precipitate? May be it works only with HNO3/H2SO4 mix?! May be NH4+ tend some intermediates to decompose?!
DNU can form a stable salts like ammonium-DNU. This salt has a poor solubility in water. It could be useful if you want to try to isolate DNU from AN/H2SO4 mix.

[Rosco Bodine]

A mixed acid density table would be handy for reference .
My guess is that the 50/50 NA/SA mixed acid of d. 1.756 would
contain SO3 to have that high of a density . For example
the 50/50 by volume average of 1.5 and 1.84 is 1.67 . And
the 50/50 by weight average of 1.5 and 1.84 is 1.65 . Unless
there is a volume decrease for the mixed acids , resulting in
an increase of density for the mixture by a "compression of solution"
type of effect .....then it is certain that the d. 1.756 mixed acid
contains SO3 to achieve that density .

I believe the only possible way to achieve your proposed reaction
using SA without SO3 , would be to form the nitrourea first . And then
you would need to use a solvent like DCM or chloroform along
with the d 1.84 SA / metal nitrate mixture , possibly using a dehydrating
nitrate salt like Mg or Al nitrate , perhaps in excess of theory , to achieve
further nitration of the nitrourea to dinitrourea . Drain cleaner SA/AN
probably wouldn't work . Any water at all would be adverse to the reaction .
Whether or not such a scheme would be workable with regards to the
hexamine component is another potential complication and unknown .

With regards to urea nitrate and nitrourea , the following excerpt is from
US5659080

In a 100 ml beaker was placed a solution of 50 grams of 33% diluted nitric acid,
and to this solution was added 14 grams of urea in small portions while stirring
vigorously. For this period of time, since the temperature does not show any rise
practically, this reaction can be carried out in room temperature and is not necessary
to be kept cool. Immediately urea nitrate resulted as crystals. Having been allowed
to stand for 10 minutes, the mixture was filtered to collect the crystals. This crystals
were washed with a little amount of water to remove the solution well and were
then dried in a desiccator under vacuum. The yield resulted in 23.8 grams of urea nitrate
(97% yield), whose melting point was measured to be 162 DEG C.

84 ml of concentrated sulfuric acid was placed in a 200 ml three-necked flask equipped
with a thermometer and a stirrer, and was cooled to -3 DEG C. To this acid was added
and dissolved 23.8 grams of urea nitrate in small portions while stirring. Having been
stirred for 30 minutes, the reaction mixture was poured on 150 grams of crushed ice.
The resulted crystals were filtered, were washed with a little amount of water for twice,
and were dried in a desiccator under vacuum. The yield resulted in 16.2 grams of nitrourea
(72.4% yield), whose melting point was measured to be 159 DEG C. Since nitrourea
dissolves easily in water, its yield is improved more by taking sufficient care when
washing with cold water.

A related patent of interest is US4513148

There was an older thread by Mr Cool which had a bit of information which
may be relevant , and I am pasting an excerpt from that old thread below

PHILOU Zrealone June 18th, 2001 06:49 AM

--------------------------------------------------------------------------------

DNU's (dinitrourea) chemistry:
Stable in H2SO4!!!
Out of a patented process:
urea is made to react with HNO3 conc (99%)and H2SO4 conc (95%) at -10C; after 1/2h the cristalls of O2N-NH-CO-NH-NO2 are forming (yield near 100% based on the urea)! The resulting cristalls must be filtered under vaccuum fitration under -10C (otherwise they redissolve) and washed with a little very cold H2O!
The remaining cristalls must be dryed; the final compound is not stable in the time due to an hydrolysis to NO, CO2 and H2O!
A stable salt can be acheived via the ammonium or potassium salt (NH4OH or KOH) but during the process 1/2 of the DNU is lost (hydrolisis)!

Now the interesting part is that when the DNU is alkylated...you get nitramines upon hydrolysis!
Thus from urea-formol polymer or from the previously mentionned methylendiacetamide or methylene diformamide...you would get the medina by simple boiling in water!
(-CO-N(NO2)-CH2-N(NO2)-)n + nH2O --> nCH2(NHNO2)2 + n CO2

------------------
"Life that deadly disease sexually transmitted".
"Chemistry is all what stinks and explode; Physic is all what never works! ;-p :-) :o)"

Mr Cool June 19th, 2001 09:16 AM

--------------------------------------------------------------------------------

Interesting...
Can you think of any reasons why the H2SO4/NH4NO3 method might not work?
If it does then I might give this a try.


[This message has been edited by Mr Cool (edited June 19, 2001).]

[VasiaPupkin]

They wrote that its mixture of anhydrous NA and SA 50:50 by mass. Without SO3. I think a volume decrease effect there.
And dichlorometane or other solvents is not necessary for it. By the way DNU is not soluble in DCM.
I have no idea why H2SO4/NH4NO3 method might not work. In theory this method have to work.

[nbk2000]

We know that it happens, in other nitrations, that the K or Na nitrate salt works better (or worse) than the NH4 nitrate, depending on the reaction.

Just try the different types of nitrate salts to find which one works the best.

[simply RED]

H2SO4 - NH4NO3 should not work in theory.
As we have a lot of unreacted NH4NO3 and H2SO4, the concentration of (NO2+) should be lower than almost anhydridous HNO3 - H2SO4 mix.
First case:
both
2NH4NO3 + H2SO4 {=} (NH4)2SO4(soluble) + 2HNO3 (HNO3 + H2SO4 {=} NO2+ + H2O + HSO4-)

Second case
only
HNO3 + H2SO4 {=} NO2+ + H2O + HSO4-

Nitrating cellulose with H2SO4 - HN4NO3 does not yeild good product.
Cellulose nitrates to pyroxilin with 95%(even lower conc) HNO3 - 98% H2SO4.

[VasiaPupkin]

I read in one book that nitration ability mixtures with H2SO4 has dependence: LiNO3>KNO3>NaNO3>NH4NO3>HNO3. Thus Nitrate/H2SO4 is even preferable than HNO3/H2SO4 from this point of view.
But in practice KHSO4 and NaHSO4 has a poor solubility in nitration mixture and its difficult to get mix. I think this problem of other kind.

[Rosco Bodine]

The nitrate plus SA reaction gives an acid sulfate byproduct for the most part
at usual temperatures using the more common nitrates like NaNO3 , KNO3 ,
or NH4NO3 .

H2SO4 + NH4NO3 -----> NH4HSO4 + HNO3

NBK is right though the only way to know if it works is to simply try the
nitration and see what results . Indeed certain nitrates may produce
different or better results than others , or all may succeed or all may fail .
There is no fixed rule to predict the outcome of such nitrations .

There have been a few past mentions made of the possibility of
making oleum from the reaction of sodium pyrosulfate and sulfuric acid ,
but no one has reported any results from experiments .

Sodium Bisulfate decomposes from heat ~400 C to form Sodium Pyrosulfate

2 NaHSO4 + Heat -----> Na2S2O7 + H2O

Na2S2O7 + 2 H2SO4 -----> H2S2O7 (liquid) + 2 NaHSO4 (solid)

Reaction conditions are unknown . This may be as simple as mixing
the two reactants and heating them together .

The liquid H2S2O7 (mp 36 C) is decanted or filtered from the NaHSO4 .

The H2S2O7 is oleum , an equimolar mixture of H2SO4 and SO3

H2SO4 + SO3 -----> H2S2O7 oleum , also called pyrosulfuric acid

This would be a very handy source for oleum if it works as simply
as it seems like it may . Sodium bisulfate is cheaply available as
a granular pH down for swimming pools , and it is also the residue
left from nitric acid manufacture from sodium nitrate .

With regards to the nitration activity of the nitrates , I have always
wanted to try magnesium nitrate and aluminum nitrate , either of
which should be easily obtainable from their sulfates and calcium nitrate ,
filtering out the calcium sulfate byproduct , and concentration of the
solution . There is actually a working range for magnesium nitrate trihydrate
to pentahydrate , where the liquid can be substituted as a dehydrating agent
for sulfuric acid , for breaking the azeotrope of 68% nitric acid and
distilling away the relatively pure nitric acid . The dehydrating properties
of the residual material can be regenerated by heating to drive off the
water and convert the material back again to the lower hydrate .

Magnesium nitrate could be effective as a dehydrating agent directly
in nitration mixtures , which could also be increased in their nitration
activity by the presence of the dissolved salt .

[Rosco Bodine]

Update

The following is from Propellants , Explosives , Pyrotechnics 26 ,
17-20 ( 2001 )

95% Sulphuric acid ( 10.9 ml , 20.0 g , 0.19 mol ) was added
to 100% nitric acid ( 13.2 ml , 20.0 g , 0.32 mol ) and then
cooled to 0 C . Urea ( 3 g , 0.05 mol ) was then added to the
reaction mixture for about 30 min . Care was taken that the
temperature of the reaction mixture wasn't allowed to rise
above 5 C . After the addition the reaction mixture was
allowed to stand for another 30 min at 0 C , during this time
a white precipitate formed . The white crystal mass was filtered
off on a glass funnel and then washed with trfluoracetic acid ,
( 5 X 5 ml portions ) . The product was dried over vacuum
giving 5.0 g ( 0.033 mol , 67 mol-% ) of N,N-dinitrourea .

The entire article was on the old forum FTP . I tried to attach
the pdf to this post but the 85K file for the three pages was
disallowed as being too large a file . There are some references
to keto-RDX in the article .

[simply RED]

Could magnesium nitrate be made completely anhydridous (easy way)?
There is a source of Mg(NO3)2 - the local chem supply. In here a kilo of it costs 2 dolars, but it is completely hydrated.

The last dinitrourea process should be tried. Using urea nitrate instead of urea could spear some good conc.HNO3.
(it should be N,N'-dinitrourea)

[VasiaPupkin]

I tried to heat a NaHSO4 a 3 years ago. It was melted during 1 hour. But I could not say something positive.
When it was mixed with 98% H2SO4 there was no smoke therefore I think Na2S2O7 was not formed. By the way an impurity of Na2SO4 is also produced when NaHSO4 treating.
Anhydrous Mg(NO3)2 could be prepared only from Mg and N2O5. Treating of pentahydrate above 90C form a dihydrate. Futher treating form a partially hydrolized product and finaly MgO.
Dihydrate is not so strong dehydrating agent as anhydrous Mg(NO3)2. And dihydrate is absolutely waste in nitrating mixtures.

[goacity]

Does anyone know any details about keto-rdx and its sensitiveness?

[Rosco Bodine]

simply Red ,

Regarding magnesium nitrate , I haven't studied all the possibilities yet ,
but it would seem perhaps that the monohydrate if it exists would be easiest
to make as a component of a melt of mixture of nitrates , where MgO is
added to a eutectic melt of anhydrous nitrates containing NH4NO3 .
I have never tried this , so I cannot be certain it would work . The
unavoidable small amount of moisture which would be in such a melt
should enable the reaction to proceed , which would be immediately
evident by an evolution of ammonia , the odor of which would prove
the reaction is proceeding and also indicate the reaction progress .

MgO + 2 NH4NO3 ------> Mg(NO3)2 + H2O + 2 NH3

It seems possible that solid chunks of magnesium scrap metal could
possibly be added to such a low temperature melt , where the
elemental magnesium might react directly , converting to magnesium
nitrate in anhydrous form going into solution in the melt with an
evolution of both hydrogen and ammonia . The magnesium scrap
might have to be activated by amalgamation with mercury for the
reaction to occur . A small amount of HgCl2 added to the melt should
work . Scrap aluminum might work in a similar way leading to a
melt containing anhydrous aluminum nitrate .

Mg + 2 NH4NO3 ------> Mg(NO3)2 + 2 NH3 + H2

Alternately to such possible schemes , heating the hydrated salts
to drive their dehydration may work to produce a degree of dehydration
which is of benefit in some nitrating mixtures , even though it may not be
possible to obtain the completely anhydrous salt by heat alone ,
without causing decomposition .

There is another idea worth mentioning with regards to magnesium nitrate .
Ordinary Epsom salt , ( MgSO4 + 7 H2O ) can be easily baked dry to the
anhydrous condition . It seems possible that addition of the anhydrous
MgSO4 to nitration mixtures containing NH4NO3 and SA , might have
a beneficial dehydrating effect directly from the drying potential of the
anhydrous MgSO4 , as well as from the presence of some MgNO3 which
may form from the reaction of the MgSO4 and the AN , according to
the ( possibly reversible ? ) reaction

MgSO4 + 2 NH4NO3 <-------> Mg(NO3)2 + (NH4)2(SO4)

This may even result in a double salt having dehydrating properties .
The literature provides no insight into this . But there is a reaction
between MgSO4 and NH4NO3 . An interesting experiment is to make
separate solutions of the two salts in the above molar proportions
and heat the separate concentrated solutions to near boiling .
Then pour the two very hot solutions together into a flexible container
which can expand without breaking , or pour the mix into an open
round bowl or basin . Immediately after mixing the boiling hot solutions
a pronounced and rapid endotherm will occur and the system may suddenly
set up to an expanding solid mass which may even produce frost on
the container . From two boiling hot solutions , you instantly get an
ice cold solid or slurry .....a very neat experiment , whatever is the resulting
material ! Anyway , the product is extremely deliquescent . I have done no
extensive study of this reaction , but it is of interest due to the uptake of water
and the thermodynamic , as well as the probable chemical compatability
with nitration mixtures where those properties may be useful .

The use of urea nitrate would not be helpful in eliminating water , since there
is no water byproduct from the formation of urea nitrate . However , there
is a molecule of water byproduct from the formation of nitrourea . So it
may be useful to form the nitrourea separately , if it can be used as a precursor
for further nitration to dinitrourea , which produces yet another molecule of water
as a byproduct . A two stage nitration would allow for use of less strong
acids in lesser amount to accomplish the formation of the dinitrourea , if the
dinitrourea does not have to be formed directly from urea itself .

The file dinitrourea.pdf ( 86 Kb ) is now in the explosives folder on the FTP .

[nbk2000]

Files larger than 50KB have to be zipped, allowing up to 500KB to be uploaded.

[simply RED]

Good info about Mg(NO3)2!
Think of urea nitrate as one part anhydridous HNO3 and one part urea.

[Jetex Kid]

MgO + 2 NH4NO3 ==> Mg(NO3)2 + H2O

kJ/mol -602 -2*366 ===> -790.7 -242

Net kJ/mol on left side of equation = -165 = No Go

[Rosco Bodine]

loop some of the H20 on the right side
of the equation back into the left side ,
or assume correctly that enough trace water
is present to kick start the reaction via

MgO + HOH ----> Mg(OH)2

and then you do have a "GO" as follows

Mg(OH)2 + 2 NH4NO3 -----> Mg(NO3)2 + 2 HOH + 2 NH3

since half the water produced is consumed making more Mg(OH)2
from the MgO , the simplified reaction may be written

MgO + 2 NH4NO3 ------> Mg(NO3)2 + HOH + 2 NH3

At too low a temperature the affinity of the nitrate for the water
might prevent the reaction , but gentle heating should drive the reaction
to the right with evolution of ammonia . Magnesium nitrate dihydrate
melts at 129 C . One potential problem could be the formation of
a basic magnesium nitrate , if one exists . It would be easy enough
to tell by weight loss and evolution of ammonia , how such a reaction
may have proceeded .

The other reaction involving the anhydrous magnesium sulfate
being added to an ammonium nitrate / sulfuric acid nitration mixture
is actually more interesting to me as an experiment .

Earlier it was mentioned the possibility of using sodium pyrosulfate
to produce fuming sulfuric acid . Of all the schemes considered for
use of a solid crystalline dessicant , which could work in a nitrate/SA
nitration mixture , sodium pyrosulfate seems to be the most likely
candidate because it either reacts with any water to revert to
sodium bisulfate , or if no water is present it may react with sulfuric
acid to produce oleum . You could regard sodium pyrosulfate as
the anhydride of sodium bisulfate , active as a chemical dehydrating
agent towards free water or towards sulfuric acid .

2 NaHSO4 + heat (~350 C) -----> Na2S2O7 + H2O

H2O + Na2S2O7 ------> 2 NaHSO4

Na2S2O7 + 2 H2SO4 ----> H2S2O7 + 2 NaHSO4

[mongo blongo]

I tried forming DNU from Nitrourea, NH4NO3 and H2SO4 and it failed again with similar results in my previous post. It's possible that the NH4+ is floating about doing naughty things in the reaction so maybe NaNO3 would be a better choice. The problem is the solubility of NaHSO4. It will be impossible to see if DNU is crystalizing out at -15°C and may cause problems when forming DHN.

[basch]

A time ago I have done some DNU. The formation was very good. A lot of white precipate apperad. I dropped then the whole acid/DNU to hexaminr/H2O. A precipate formed and i filterd it out.
I put it outside to dry it at the air. But a few days later it was nearly all gone?!!
I thnik it has been hydrolysesd by water or it is very hygroscopic so that in has been dissolved in water. But I am noit sure. What do you think.
ps I'm sorry for my bad english!

basch