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Axt
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[*] posted on 18-3-2005 at 04:57
Nitrated Organic Acids & Salts


Can't report any successes here, but someone may find a useable route to a useful exposive product.

Nitration of organic acids contaning hydroxyl groups, such as citric acid will result in nitric esters that have the ability to form salts. These may be useful OTC primary explosives. Citric would go to a mononitrate, but the acid that would be of most interest, and relatively easily available is tartaric acid, nitrated to its dinitrate would result in a 0%OB explosive (though, admittedly largely bound up in carboxyl groups).



I've only seen somewhat vague references to sodium and calcium salts of tartaric acid dinitrate, but one would think addition of Pb acetate to a solution of the acid nitrate would result in the precipitation of 'Lead dinitroxytartrate'. This isnt as simple as it seems!

Tartaric acid dinitrate (TADN) is readily made by the addition of 200ml 98% H2SO4 to 25g tartaric acid in 50g 70% HNO3. Added at 15°C with no further cooling the temperature will reach 45-50°C on addition of the H2SO4, this is then left to cool slowly over the course of a couple hours at which point it results in a solid paste of crystals as the dinitrate precipitates.

Heres the problem, isolating the free acid from the nitration mixture is very troublesome (TADN is soluble in water), ideally one would use vacuum filtration to draw off most of the viscous liquid, then add a minimum of water and extract with ether. Thats no OTC method. And ether can't be added directly to the nitration acids, so good filtration of the strong viscous acids IS necessary.

I have tried neutralising the whole lot with bicard and precipitating with copper sulphate (since the sulphate is soluble), but no product. Pressing as much liquid from the crude TADN then precipitating with Pb-acetate only resulted in Pb-sulphate, as it holds onto the H2SO4 very well. Attempts to flush the H2SO4 out with DCM failed also.

Anyway, what I thought was a good idea at the time really wasn't realised, so someone else may try to get a useful salt, whichever way possible. The free acid is explosive, but sensitive and unstable towards water (and hygroscopic!).

Nitration of Malic acid (which is tartaric minus one hydroxyl group) was also attempted, but doesn't result in a precipitate from the nitration acids, it may be a liquid mononitrate.

It may be worth trying the nitration with strong fuming HNO3, which is more readily extracted then H2SO4, and wont form insoluble precipitates from Ag/Pb salts.

Citric, tartaric and malic acids were purchased from a food supply store.

[Edited on 9-12-2005 by Axt]
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[*] posted on 18-3-2005 at 21:13


Did some searching today at school after reading this topic, from Journal of the American Chemical Sociotey, Wolfrom, M.L., Rosenthal, Alex. Vol 75 [3662] Nitrated Aldonic Acids
Quote:
[Absract]
D-Gluconamide and D-galactonamide have been nitrated by the method of Caesar and Goldgrank to yield the pentanitrates which on deamination with nitrosyl chloride produced the pentanitrates of D-gluconic and D-galactonic acids. The methyl esters of the later were obtained with diazomethane. These nitrated amides and methyl esters were compatable with cellulose nitrate. The compoundes were explosive and were unstable toward moisture.
Interesting article overall but specifially the part that caught my eye was in the first paragraph:
Quote:
Nitrates of hydroxyl acids and their derivatives have been little investigates. The liquid lactic acid nitrate [2,3] and the crystaline L-(+)-tartaric acid dinitrates [4-6] have long been known; they are unstable and undergo interesting and well established transformations. Duval [7] prepared glycolic acid nitrate, and glyceric acid dinitrate in crystaline form. Isoamyl lactate nitrate [8] and the nitrates of methyl [3], propyl [8] and butyl [8] glycolates have been reported as liquids. It ws the objective of the work herein recorded to prepare the pentanitrate of aldonic acid. It was considered that the carboxylate ester of such a substance might be a suitable explosive plasticizer for cellulose nitrate.


[2] L. Henry, Ber., 12, 1837 (1879).
[3] H. Duval, Compt. rend., 137, 1262 (1903); Bull. soc. chim., {3} 31, 243 (1904).
[4] V. Dessaignes, Compt. rend., 34, 731 (1852)
[5] A. Kekule, Ann., 221, 230 (1883)
[6] A. Lachman, THIS JOURNAL, 48, 577 (1921)
[7] H. Duval, Compt. rend., 137, 571 (1903); Bull. soc. chim., {3} 29, 601 (1903)
[8] S.E. Forman, C.J. Carr and J.C. Krantz, Jr., J. Am. Pharm. Assoc., Scientific Edition, 30, 132 (1941)

Attached is reference [6] on nitro tartaric acid. It contains the preparation of nitro tartaric acid and basically restates the information Axt said above on the properties of nitrotartaric acid.

[Edited on 3/19/2005 by BromicAcid]

Attachment: nitrotart.pdf (336kB)
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[*] posted on 21-3-2005 at 08:07


So would making a nitrated aldaric acid go something like

aldose sugar (or any reducing sugar) --dilute nitric acid--> aldaric acid --mixed acid--> nitroaldaric acid ?

And wouldn't lactic acid be a good nitration substrate also?

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[*] posted on 21-3-2005 at 15:14


Can you elaborate more on the process of oxidation with dilute HNO3? for example the nitration of glucose results in a pentanitrate, but the aldehyde group is preserved.

glucose --> glucuronic/glucaric/gluconic/whatever acid via dilute HNO3 would be interesting and readily available.

I cant say if any of the acids will produce a useful explosive, I was only interested in the possibility of metal salts. Though one would think if forming salts were possible I would be able to find references to them, but nope. It could be that they are no more stable then the free acid itself.

Bromic, the article you attached is the one I derived the syth from, where there is mention of the sodium/calcium salts, though no mention of their stability/explosablity. I just halved the amount, and converted volume measurements to grams.
There another synth for tartaric acid dinitrate on orgsyn here as a precursor to imidazole. And it is mentioned in PATR, though only to say "The dinitrate is considered too sensitive for use in a military explosive". No properties/synth or mention of its salts.

[Edited on 21-3-2005 by Axt]
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[*] posted on 23-3-2005 at 03:31


Axt, sorry, I can't find my notes on how to perform the oxidation, but I do remember that glucaric acid/saccharic acid is produced by heating glucose with dilute nitric acid.

Gluconic acid is the oxidation product of glucose using milder oxidizing agents; glucuronic acid is produced by vertebrates to conjugate ingested toxins. Hope this clears stuff up.

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[*] posted on 24-3-2005 at 17:25
Nitration products of Amides of tartaric/citric/oxalic acid


As an idea:
In Ruldoph & Meyer, Explosivstoffe, it mentions dinitrodimethyloxamide, or the related compound dinitrodioxy ethyloxamid dinitrate (page 96 both).
Now don't let that confuse you, essentially it is the amide of oxalic acid that is nitrated, i.e. in the first case it is oxalic acid reacting with methylamine, forming dimethyloxamide, and in the second case it's oxalic acid reacting with ethanolamine rather than methylamine, forming diethanol oxamid.

Both can be nitrated with the mixed acids HNO3/H2SO4, so presumably massive HNO3 concentrations aren't required.
Nitration occurs on the amide nitrogen, and in the latter case, also on the OH from the former ethanolamine.

Now the point is, I imagine, the same could be done with tartaric acid - react it with methylamine, forming (CH3NH)CO-CH2OH-CH2OH-CO(HN-CH3), which can be nitrated on both the hydroxyls AND the nitrogen, forming a presumably more energetic material than tartaric acid dinitrate alone.
The product would be
H3C-N(NO2)-CO-CH2NO3-CH2NO3-CO-N(NO2)CH3

Further, of course, this can be done with ethanolamine.

I was looking for publications that described the making of the oxamides. Hopefully these conditions can be employed for generating the amides of tartaric acid, or even citric acid.
The nitration after that should be rel. easy!

So no worries Axt, you aren't going to be unemployed anytime soon ! :D

[Edited on 25-3-2005 by chemoleo]




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[*] posted on 24-3-2005 at 19:01


Pg 103 in english free version of "Explosives".

Did you realise that was the topic of the patent I mistakenly sent you?
Does it give any indication of the reaction conditions? VOD and other properties are given PATR2700, D 1255. VOD = 7860m/s @ 1.65g/cm.

Using the same nomenclature I guess you would call the tartaric acid derivative as "dinitrodioxyethyltartramide tetranitrate" C8H10N8O18, OB=-9.48%. No references to it anywhere though.
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[*] posted on 24-3-2005 at 19:43
Addendum


Yah, I realised that, but I meant to make the point not so much of nitramide formation (although admittedly, upon reconsidering it, this is a decisive added bonus:D), but rather of derivatives of hydroxylated organic acids that can be nitrated (which are likely to be barely soluble in H2O, judging by the trends stated in the patent/R&M - i.e. the nitroxamides are insoluble, and i.e. NG is also insoluble, plus there aren't any free charges (such as with the tartaric acid dinitrate where the free COOH causes its water solubility)).
Hence this may provide a route for purification, which seemed not easy for the underivatised tartaric acid dinitrate (which is water soluble as you say).
Sorry if you felt you didn't get credit or something for sending me the patent..... many thanks, I'd likely not write this without having read it .. that's what information sharing and many minds thinking about the same thing is about :)

So essentially, the point was that organic acids such as tartaric/citric acid, or even glucuronic (etc) acids could be used to form amides, which then can be nitrated to form both nitrates and nitramides (within the same molecule), something I believe wasn't said in the patent or anywhere else. That was essentially the point of my previous post.

Admittedly, this is all hypothetical at this point - they may not exist, but at looking at the chemistry I can't see why an oxalic acid nitramide would exist, but not a tartaric acid one, considering that tartaric acid can be nitrated on its own, and equally dimethyl oxamide to form the nitr(ox)amide.

Anyway, don't you think it might be worth trying it?




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[*] posted on 24-3-2005 at 20:20


Quote:
Originally posted by chemoleo
Sorry if you felt you didn't get credit or something for sending me the patent..... many thanks, I'd likely not write this without having read it .. that's what information sharing and many minds thinking about the same thing is about :)


That wasn't what I was after at all. Because you didn't refer to the patent I thought you may have not made the connection, thus also missing any mention of the reaction conditions relating to condensing oxalic with ethanolamine. Which I can now assume are not mentioned in it. (We are talking about GE543174).

Anyway, I believe you are the one with ethanolamine on hand, are you not ;) It would be interesting if it did work as it would open up routes to a large range of nitramide-nitrates. But But But! it seems oxalic acid wont condense directly with amines, but needs an oxalate ester. Dibutyl oxolate is given in PATR.

Though this seems sentence may refer to a more direct route "This compd was first prepd by Von Herz (Ref 2) by condensing ethanolamine with oxalic acid, and nitrating the product.". Ref 2 is the patent. Thats why I had to bring it up.

Oh, and I gave wrong pg, D1244 vol.5.

[Edited on 25-3-2005 by Axt]
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[*] posted on 24-3-2005 at 20:42


Ok, fair enough. Yes, that patent did not mention how the oxalic acid diamide was made.
They made it sound simple though.

Ah. I see your edit. It needs esters. That's interesting.
So R-CO-OR + H2N-CR3 --> R-CO-HN-CR3 + HOR.
I wonder why the free acid wouldn't work. I am sure one of the budding chemists out there has an answer for this?

Yes, you are correct. Ref 2, in PATR page D 1244 says it can be done by direct condensation of acid with amid.
Makes kinda sense, as amino acids can also be condensed directly to peptides. As a result I can also understand why proteins, when un-hydrolysed, are good nitration targets - the NH from the amid bond (NH-CO) forms N(NO2)-CO).
. Time to dig up that reference!

[Edited on 25-3-2005 by chemoleo]




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[*] posted on 27-3-2005 at 09:19


The amine and the acid will neutralize each other first! :D

Anyway, IIRC, amines are not nucleophilic enough to push away the -OH group attached to the carbonyl.

Is someone suggesting nitrating amino acids already? How about glutamic acid and glycine for starters?

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[*] posted on 28-3-2005 at 08:33


I successfully nitrated glycine. Glycine nitrate looks similar to HDN but the crystals are smaller. I havent detonated it yet. I think its fairly insensitive. But there is not that much information about this "explosive".
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[*] posted on 28-3-2005 at 08:57


I too have Nitrated Glycine, but it is extremely exspensive around me so it was a two time deal. There is actually alot of info about Glycine Nitrate, in fact there is a thread here;
http://67.15.145.24/~sciencem/talk/viewthread.php?tid=2642

And the references from PATR2700 for Nitroglycine and Glycine Nitrate are in Volume1, A 178.

Edit:I realize this does not fit into this thread but I was trying to keep the thread on track by telling Kinpack where to go to discuss Glycine Nitrate, no worries;)


[Edited on 28-3-2005 by Joeychemist]
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[*] posted on 28-3-2005 at 09:02


Kinepak/Joeychemist, your 'nitrated glycine' is just glycine nitrate, an amine salt, it isn't a 'nitrate' as in glycerol trinitrate, hence it doesn't really fit into this thread. Please stay on topic.
Nitroglycine exists too but can't be made that way. There's a thread on this somewhere. Again not really on topic.

Sparkgap, I don't understand.

R-CO-NH- R can be nitrated, to give the nitramide, i.e. R-CO-N(NO2)-R. There are plenty examples. Amide bonds don't form salts easily (i.e. R-CO-(NH2+)-R, because of its resonant stabilisation with the neighbouring CO, that's why amide bonds are so stable.


Nitrating amino acids - whats the point? All you get is the respective salts, which have a terrible OB except glycine. Serine and threonine may be an option however, as they are hydroxy amino acids that could be esterified with HNO3, while forming the nitrate salt with the NH2 at the same time.
Similar to ethanolamine forming the nitrate ester and nitrate salt at the same time, [H3N-CH2CH2-ONO2][NO3]

The point of all the posts above was that

1) organic food acid such as tartaric acid can be amidated with i.e. methylamine, to form the diamide.
2) This can be nitrated, not just on the free OH of the organic acid, but also on the amide bond
3) the resulting product is likely to have a bad solubility in water (unlike the free tartaric acid dinitrate), which is good for purification, and thereby avoids the problems of normal organic acid nitrates.


[Edited on 28-3-2005 by chemoleo]




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[*] posted on 29-3-2005 at 02:54


chemoleo, I am now confused. I was under the impression that your question "I wonder why the free acid wouldn't work?" was in reference to preparing the amide from acid and amine. My answer is valid if so. Nitrating the amide, however, that looks OK. Mixed acid seems to be required for this, however. This reminds me to post the rest of the "Nitration" book!

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[*] posted on 13-1-2006 at 20:14


Another thought occurred to me after reading the GHB thread - There, GABA is diazotised (sandmeyer), and gamma hydroxy butyric acid (after hydrolysis of the lactone) is obtained.

A similar treatment can be done with a varitey of amino acids, which are all available in health stores:

Glycine NH2-CH2-COOH --> Sandmeyer --> 2 hydroxy ethanoic acid, HO-CH2-COOH
Serine CH2OH-CH(NH2)-COOH --> 2,3 hydroxy propanoic acid, CH2OH-CHOH-COOH
phenylanaline Phe-CH2-CH(NH2)COOH --> Phe-CH2-CHOH-COOH


I should think that particularly the former two would be an interesting and novel target for nitration. Something else than just the widely available citric/tartaric/maleic acids!

[Edited on 14-1-2006 by chemoleo]




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[*] posted on 3-12-2006 at 14:57


Excuse me terribly for bumping an ancient thread.

If I were to introduce an amide into an acidic environment, would the N or the O be protonated? Due to the resonance-induced stabilization I would expect the O to be accesible to protonation but not the N.

And is it safe to say that the N is more like sp2 hybridized than it is like an sp3 hybridized particle? After all it uses the p orbital for the resonance stabilization.




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[*] posted on 5-12-2006 at 10:33


Possible solution for isolating free acid? Prehaps you could try making an ester out of the nitrated organic acid. The ester could possibly be easier to isolate or just as difficult to isolate.



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[*] posted on 27-6-2016 at 14:49


Heh, 10 years ago someone apologized for bumping a one year old thread. Now it's ~10 years old...

I was wondering about the nitrocitric acid. I'd poked around and found Journal of the Chemical Society Vol. 29 contains a short reference to nitrocitric acid:

Quote:
CITRIC acid is dried till one molecule of water is evolved, and is then introduced into a mixture of 1 part of colorless fuming nitric acid and 2 parts of pure sulphuric acid. Some heat is evolved While the citric acid is dissolving; on standing for some days nitrocitric acid crystallizes out partially. It is transformed into the barium Salt, and after separation of insoluble barium sulphate, is converted Into [...] Nitrocitric acid is insoluble in ether, but dissolves readily in alcohol. It appears to give an ether when treated with hydrochloric acid and alcohol. Analyses of the lead and barium salts showed the formula be [...] Nitrostearic acid [...], prepared in a Similar manner is a pale yellow powder soluble in alcohol.


Anyone tried this? Do I have much chance making it with ammonium nitrate, 95% conc. sulfuric acid, and citric acid?
Thanks, sorry for bumping this.




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[*] posted on 28-6-2016 at 02:22


The use of NH4NO3 will complicate the process (very viscous because of microcristal formation) and the isolation since you will get also other possible precipitate/crystals:
-NH4HSO4
-(NH4)2SO4
-NH4NO3
and maybe
-monoammonium citric acid nitrate
-diammonium citric acid nitrate
-triammonium citric acid nitrate
aside from the citric acid nitrate

Better then mix your NH4NO3/H2SO4, allow to heat to fully dissolve; then cool for crystallization of most of the NH4HSO4/(NH4)2SO4, then take the supernatant (rich at concentrated HNO3 with traces of NH4 salts and H2SO4) for your nitration.

About the tread compounds:
1) The unstability of hydroxy acid nitrates esters comes from the inherent acidity and solubility.

2) The hydroxy acids are more acidic than the related acid without the hydroxy (about 10 fold --> 1 pKa unit).

3) Alkylic nitric esters don't like acidity traces...this is even more true if the acid comes from the same molecule...

4) The carboxylic moeity cares for water solubility what is bad for nitric esters because it favourizes hydrolysis.

5) The presence of the C=O of the carboxylic allows for resonance structure and proton jump:
R-CH(-ONO2)-C(=O)-OH <==--> R-C(-ONO2)=C(-OH)2
Even if minor this compound will be very pleased to exchange its NO2(+) for a H(+) from water (hydrolyse) thus setting HO-NO2 free.

6) Because of the easy HONO2 release, the use of strong bases is precluded for salt formation (tartric acid dinitrate is thus a potential 4 equivalent acid and will be neutralized not by 2 NaOH, but by 4...)

7) Because of the easy HONO2 release, many solvents are uncompatible into the process of extraction-recrystallization...
Aceton and Isopropanol are uncompatible with HNO3; alcohols will be involved in partial transesterification...
R-CH(-ONO2)-CO2H + CH3-OH <==> R-CH(-OH)-CO2H + CH3-ONO2

Thus all this must be a severe warning that those stuffs are not to play with in large quantities.

I have wel some ideas to circumvent the troubles encountered by previous operators in making salts....
1°) Use HNO3 conc and eventually Ac2O (will make eventually mixed anhydride Ac-tartric, Ac-nitrate) and as a bonus all acetates salts are soluble just like nitrates.

2°) Use carbonate salts instead of hydroxydes...carboxylic acid all displace carbonic acid from its salts leading to CO2 bubbling; this is even more true for hydroxycarboxy acids and their esters.
Also most unsoluble carbonate are neutral and so they will not favourise hydrolysis --> CaCO3, PbCO3, ...
Alkaline carbonates may be too basic (Na2CO3 for example) so better use hydrogénocarbonates (NaHCO3).
NH4HCO3 and (NH4)2CO3 should be just fine.

3°) Instead of starting from tartric acid (dihydroxysuccinic acid/dihydroxybutandioic acid) and HNO3, one may think to start from the halogenated compound maleic acid dihalide (dihalobutandioic acid) or fumaric acid dihalide (dihalobutandioic acid)...Maleic acid is cis-butendioic acid and fumaric acid is trans-butendioic acid and they will add dihalogen very easily on the double bond.
Once there you get alfa halogenated carboxy acid...
HO2C-CH=CH-CO2H + Br2 --> HO2C-CHBr-CHBr-CO2H
HO2C-CHOH-CHOH-CO2H + HBr + catalyst -->HO2C-CHBr-CHBr-CO2H
Such alfa halogenated carboxylic acid do display a high reactivity of the halogen because of the electron withdrawing effect of the viccinal CO2H...this means they are easily exchangeable...
Alkyl halides do react with AgNO3... primary > secondary > tertiary...here it is secondary but activated....
HO2C-CHBr-CHBr-CO2H + 2 AgONO2 --> HO2C-CHONO2-CHONO2-CO2H + 2 AgBr(s)(precipitate)
HO2C-CHONO2-CHONO2-CO2H + AgNO3 (excess) --> AgO2C-CHONO2-CHONO2-CO2Ag (precipitate?) + HNO3


[Edited on 29-6-2016 by PHILOU Zrealone]




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