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Author: Subject: Dinitroresorcinol preparation
roXefeller
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[*] posted on 14-2-2018 at 06:37
Dinitroresorcinol preparation


So a funny thing happened on the way to the forum. I was studying patent 4745232 in order to prepare dinitroresorcinol (or mono, but not trinitroresorcinol). In this the author posits that the presence of the nitrosonium ion contributes to the nitro substitution of the 1- position of resorcinol or resorcinol derivative. As a fix for this, an addition of urea or hydrazine or other reagent is added to oxidize the nitrous acid leaving mostly nitronium ions present for the reaction which will preferentially substitute to the 4,6- positions. The attached article experimentally concludes that the acidic pathway of urea and nitrous acid leads to oxidation but also ammonium ions. That's the background at least. Two days ago I performed a nitration following the patent description. 5.2g of HNO3 soln (70%, reagent grade) was combined with 0.5g urea. 16.2g water was combined with 86.6g H2SO4 soln (96%). I combined the acids and let them cool to 0C. Some bubbling was occurring while this was happening. After they cooled I put it on a stirrer with ice bath and added by drops 5g resorcinyl diacetate (as far as I know, prepared from resorcinol, NaOH, and acetic anhydride). The nitration proceeded much like a phenol nitration. Yellow precipitate started to form, entrained in the liquor. Eventually the liquor darkened and the yellow precipitate showed brightly. The patent recommends a 3 hour reaction time, so I left it to stir (like with phenol), but after no more than an hour the precipitate had dissolved completely into the liquor. The new solution was a intensely dark blue. That was new, not something I see with phenol. I knew something had changed so I dumped the beaker into cold water to stop the reaction. Precipitate formed and the water turned red. I haven't seen blue since, so I think it is a pH indicator which shows blue for low pH. I left the crash to settle overnight. I decanted the top portion and filtered the precipitate out, disposing of the wash water. I left the solids to dry, mauve in color. When mostly dry (a little clay-like) I poured in ethyl acetate which immediately dissolved the red color leaving a tan solid. I filtered the solids, reserving the liquid. In a watch glass, one mL of the red solution evaporated to some reflective crystals. The solids are still tan. Neither show a propensity for flame decomposition that I would expect from the nitration product. They melt and then burn slowly leaving black crud.

I was looking up what resorcinol derivative turns blue, and I saw some data that leads to an amino substitution, which maybe comes from the decomposition of urea with nitrous acid. Admittedly, I added 5x more urea than the patent example, but well within the specification of 0.01%-10% that it claims. I didn't take a smell of the product to see if it smells of amine. I'm also hoping still to get melt points of both. Could the ammonium ion also act as a catalyst to condense two dinitro products into a polybenzene compound?

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byko3y
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[*] posted on 14-2-2018 at 18:54


Blue is the color of indophenol, qualitative reagent made of sulfuric acid and sodium nitrite is called libermann reagent. Important condition for indophenol formation is a large excess of phenol relative to nitrite ion.

[Edited on 15-2-2018 by byko3y]
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Boffis
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[*] posted on 15-2-2018 at 00:45


The only problem with the Lieberman reaction idea is that it requires nitrous acid or nitrite proxy to be present and the whole idea of the urea is to specifically remove these. The fact that you can isolate a crystalline red material suggests that quite a lot was produced.

Under basic conditions ammonia reacts with resorcinol and atmospheric oxygen to give a litmus like product that was at one time used as a pH indicator under the name resorcein if I recall correctly. However, here you seem to have strongly acid conditions and the ammonium ion concentration must be very small so I would that rules out this possibility too. I also find it hard to believe that you could get reduction of the nitro groups under these conditions too.

Have you tried testing the red crystals to see if they have reversible pH sensitivity?
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roXefeller
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[*] posted on 15-2-2018 at 06:39


The red substance in the ethyl acetate isn't reacting to pH. I've put it in 50% H2SO4 and again in NaOH solution without change.

The tan solid has a melt point around 240C. The ethyl acetate solution has grown some mauve crystals with melt point at 225C. They smell like lipstick. Not sure what that contains.
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byko3y
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[*] posted on 15-2-2018 at 06:41


It's physically impossible to destroy 5.2 g of 70% nitric acid with 0.5 g of urea. Something is left behind and reacts. Indophenol is a strong dye.
Yes, the resorcinol can be oxidized with ammonia and air ( Über Orceinfarbstoffe, XVIII. Die Autoxydationsprodukte des Resorcins und 2-Methyl-resorcins in Ammoniak, Chem. Ber. 96(6), p. 1579), but the condition we discuss is not basic.
There's also some qualitative test performed by reacting phenol with acidic nitrite and right after the mixture boiled it is added to a solution of ammonium chloride (New specific tests for distinguishing carbolic acid, the cresols and certain other phenols, Analyst, 1927,52, 335), but I'm not sure about its mechanism.
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roXefeller
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[*] posted on 15-2-2018 at 09:44


I've repeated the workup again with two differences. First I used 0.12g urea. Second I combine the acids first and cool them. Just before beginning the nitration is when I added the urea. Previously the urea sat with the nitric acid for several minutes before combining with sulfuric. Now once it dissolved I began resorcinol diacetate additions. The liquor didn't bubble as before and didn't have a prolonged time with the urea while cooling. It's been nearly three hours now and just looks like a milky light brown suspension.

[Edited on 15-2-2018 by roXefeller]
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[*] posted on 16-2-2018 at 04:19


I've got a theory question if you could answer. With the presence of the acetate ester acting to reduce the electron resonance by 50%, the nitration proceeds slower than with the non-ester diol. So there is for a time a 4,6 dinitroresorcinyl 1,3 diacetate? What causes the ester to break and return to a diol group? I find in the reaction (the one that doesn't turn blue) that solids form almost immediately, a minute at most. And for the rest of the time they are merely held in suspension assumably nitrating. What in theory is this intermediate? Will the addition of NaOH react with the ester and result in sodium acetate and dinitroresorcinol as it would with ethyl acetate?
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[*] posted on 22-2-2018 at 01:10


@byko3y, That paper from the Analyst is very interesting and is certainly enigmatic but here the nitrite is used in conjunction with nitrate and then added to the ammonia which being in excess generates basic conditions. In the OP description he has added urea which should have removed the nitrite ions but introduced ammonium ion, though still under acid condition. I will have to try this test an see if other bases can replace ammonia (ie its a pH effect) or whether its specific to ammonia.
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roXefeller
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[*] posted on 26-2-2018 at 14:28


After testing some mg quantities of the tan precipitate, I do believe that it is dinitroresorcinol diacetate. In water in a test tube a small quantity will not dissolve except when adding a small amount of NaOH. In very small amounts the red diol forms as a precipitate. With another addition of NaOH the sodium salt forms which is readily soluble. The addition of HCl causes a tan precipitate again, so maybe the acetate is reformed? I tried a larger workup to test this. 1560g of 96% H2SO4 was diluted with 291g water and cooled in an ice bath. When cooled to 12C, 92g 70% HNO3 was added. When all was ready, 1.8g urea was added and mixed in. 90g resorcinol diacetate was added with stirring in 10ml quantities maintaining the temperature below 20C. The reaction was left to proceed for another 3 hours.

Portions of the mixture were poured into ice water and vacuum filtered until all precipitate was collected. The tan diacetate has a real tendency to clump when damp causing the final drying to be time consuming. Vacuum desiccation may be a good option. I collected 192g of the rinsed, wet product. 69g when dry, 59% yield, likely due to high temperatures.

I attempted to determine solubility of this. It will dissolve in methanol, but not well. Sigma-aldrich claims that the dinitroresorcinol (the diol) is "transparent" in methanol. I'm not sure what this means, insoluble? At first I assumed that this meant it was fairly soluble, so I added 61g of the diacetate to warm methanol, but even with 1200ml it didn't dissolve sufficiently. So I proceeded to add 18g NaOH in 100ml methanol to produce the diol. However it never dissolved after reacting. Worried that due to the undissolved nature, half of the product was unreacted and half went all the way to the sodium salt (sodium dinitroresorcinate), I added another equal quantity of NaOH in water to ensure enough sodium was present for a full conversion to the sodium salt. Never was it very soluble, which disappointed me. I had hoped that it would dissolve in the methanol and I could recrystallize with a dropwise addition of cold water. The goal would be to have the diol form in large crystals on the bottom of a mostly water solution with the sodium acetate remaining in solution. Instead I have around 500ml of red methanol solution (presumably with most of the sodium acetate), and 49g of orange precipitate that also resists drying.

The sodium salt is very water soluble. I put a couple large pieces in a small test tube before it wouldn't hold more. When I added lead nitrate solution as an analytical test, I got an immediate bright orange precipitate, which upon drying exhibits the expected thermal sensitivity. So with the solubility and this salt reaction, I'm fairly confident that it is mostly sodium salt. I don't know how much sodium acetate is contaminating it. If there is any, it will participate in the lead nitrate reaction just the same. I suppose at this point I could dissolve some in water and acidify to the diol, which should precipitate for collection. Then I could determine what its solubility is (which isn't really reported anywhere). Are there any solubility models that would predict this molecule? What I worry about is if any sodium acetate is present in solution, would I get dinitroresorcinol diacetate again like I was seeing in my test tube experiment. If I knew the pure melting point of the diol maybe I could decide if sodium acetate is contaminating.
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Boffis
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[*] posted on 27-2-2018 at 01:42


roXefellor, have you tried recrystallizing a little of the diol from boiling water? I would expect the free diol to be much more soluble in hot water and sparingly soluble in the cold. The diacetate should be much less soluble. Once hydrolyzed the acetate will not reform in diluted solutions, the precipitate on the addition of HCl is almost certainly the diol liberated from the Na salt.

To check whether the tan material is the diacetate or the diol I would try treating a little of it with acetic anhydride again and checking either the melting points or the solubilities of starting material and product.

In the patent you referred to above it is pretty clear that the product of nitration on resorcinol diacetate is the free dinitro-diol so that the acetate must undergo hydrolysis. Why do you feel that this might not be the case and you are getting the dinitro-diacetate?
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roXefeller
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[*] posted on 27-2-2018 at 05:21



Quote:

Once hydrolyzed the acetate will not reform in diluted solutions, the precipitate on the addition of HCl is almost certainly the diol liberated from the Na salt.


I thought this too, but the tan precipitate seemed too coincidental. The crystals were reddish orange in the test tube and very insoluble when I added only a slight amount of NaOH, so I presumed that was the diol color. I further supposed that the activating groups on the ring maybe made the ester more favorable when HCl was added. But beyond the few grains in the in test tube, I haven't isolated any free diol for further testing.


Quote:

In the patent you referred to above it is pretty clear that the product of nitration on resorcinol diacetate is the free dinitro-diol so that the acetate must undergo hydrolysis. Why do you feel that this might not be the case and you are getting the dinitro-diacetate?


It does read that way, but patents are also written for "those knowledgeable in the arts" so they aren't always forthcoming about details that the author deems omittable. I think if the acid conditions were concentrated or the temperatures higher, the acetate would hydrolyze but it would also promote more thorough nitration which was counter to the objective. Additionally, acid catalyzed hydrolysis is an equilibrium reaction (as opposed to base catalyzed) so the hydrolysis might be present but not complete. My product also doesn't burn well, which TCI Chemicals (and others) indicate in the MSDS listing. Notably, TCI provides a 215-218C melting point that I could compare to. It would be nice to get an empirical C:H:O:N combustion ratio but I don't think I have easy access to gas detection equipment. But I will try melt point comparisons with the diacetate, the base hydrolyzed diol, and the acidified sodium salt.

[Edited on 27-2-2018 by roXefeller]
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[*] posted on 1-3-2018 at 04:00


I ran a test last night. 121mg of what I think is sodium salt was dissolved completely in water, water was orange. I added HCL dropwise until no further changes occured. Water was clear, precipitate was faint yellow. I left this to dry. Taring the scale with a second filter I weighed 92mg of precipitate. Theoretical weight is 99.2mg. I don't have enough filters to measure their average weight and standard deviation, but 7mg seems like a reasonable discrepancy. I loaded this into a capillary tube. I found it melted at 217C. TCI chemicals lists it at 215-218C. I'll spend a little time determining solubility of this free diol before I test what I believe to be the diacetate and the free diol produced from that.
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[*] posted on 1-3-2018 at 07:09


@roXefeller, that sound pretty promising, it will be interesting to see what melting point you get for your "diacetate".
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[*] posted on 1-3-2018 at 07:56


Knowing how hydrogen bonds affect physical melting mechanisms, how would you predict the change in melting point between a diol with available hydrogen and an acetate ester? Ethanol has a melting point of -114C and ethyl acetate has MP of -83C, so the extra group between the carbonyl and methyl cause it to freeze at warmer temperatures. For phenol the MP is 40C and phenyl acetate the MP is -30C. The same extra carbonyl and methyl groups cause a reduced ability to freeze at warmer temperatures.

Studying the effects of hydrogen bonds and dipole attractions, I would say that the diol has an intramolecular hydrogen bond that lowers the melting point and decreases solubility much like o-nitrophenol. Esterifying the two hydroxy groups to acetate should have the effect of increasing melting temperature because it replaces the ineffective hydrogen bonds with oxygen lone pair dipoles. That's my supposition at least. As for the 69g batch of what I think is diacetate, I think it's a mix. It definitely begins to melt around 217C but I think there is a component in there that isn't fully liquifying until higher. I feel like I need to try it in an oil bath with a test tube. Certainly at least 25% is hydrolyzed in the nitration.

I also reacted the sodium salt with acetic anhydride. It eventually reacted and settled out. Though it has a really low melting point, like 130C.

[Edited on 2-3-2018 by roXefeller]
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