Sciencemadness Discussion Board - HNO3 is a terrible NO3- source...why use it?

HNO3 is a terrible NO3- source...why use it?

niertap - 12-3-2012 at 01:58

In most energetic material synthesis' 70% HNO3 is used. Why doesn't everyone just use KNO3 or NaNO3? There will be a considerable amount of water that will not have to be removed from the reaction, decreasing the amount of H2SO4 needed.

A 3.8L jug of 98% H2SO4 can be purchased from the hardware store for around $25. Why would someone not in an analytical setting use anything different?

Bhaskar - 12-3-2012 at 02:07

Hi,
Your suggestion is quite right, I quite agree but how can the acidification of a nitrate salt form nitric acid above it's aezotrope? I don't know myself so as your suggestion says, you must be knowing the maximum concentration reached by a nitrate salt on acidification, say sulfuric acid in nitrations?

Bot0nist - 12-3-2012 at 04:53

Nitrations are not a "one size fits all" reaction, and are tailor fitted to the substrate. In some nitrations sulfuric acid can not be used at all, like hexamine.

Bhaskar - 12-3-2012 at 06:20

Quote: Originally posted by Bot0nist

Nitrations are not a "one size fits all" reaction, and are tailor fitted to the substrate. In some nitrations sulfuric acid can not be used at all, like hexamine.

He's just asking wether he can use nitrate salts for performing nitrations, how is what you just said the answer to the question?

barley81 - 12-3-2012 at 06:31

The answer is yes. Search for the syntheses of nitrocellulose and picric acid on this forum and you will find that some use nitrate salts.

Bot0nist - 12-3-2012 at 06:58

Quote: Originally posted by Bhaskar

Quote: Originally posted by Bot0nist

Nitrations are not a "one size fits all" reaction, and are tailor fitted to the substrate. In some nitrations sulfuric acid can not be used at all, like hexamine.

He's just asking wether he can use nitrate salts for performing nitrations, how is what you just said the answer to the question?

No, he asked why is nitric acid used at all. I assumed that the OP was not completely dense and was capable of utfse. If he is capable of such a seemingly difficult feat for most, then he would already know the answer to the question you thought he asked. I, however, was attempting to answer the question he actually asked. I Think.

How could we employ nitrate salts in a nitration without dissolving then in concentrated sulfuric acid?

What if there is issues with you isolating your nitrated product due to the contaminates (potassium sulfate, etc) you added?

The easiest answer to these two questions is to distill nitric acid from the nitrate salt and sulfuric acid. No added contaminates during the nitration, and no need to use sulfuric, like when it's not compatible with the substrate, intermediate, or target.

[Edited on 12-3-2012 by Bot0nist]

Bhaskar - 12-3-2012 at 08:54

I see... :cool:

Pulverulescent - 12-3-2012 at 09:10

If one is unable, or unwilling (too lazy) to prepare strong HNO3 by distillation then nitration by H2SO4/nitrate-salt mixtures would seem to be a doable alternative, poor though it is . . .

Bhaskar - 12-3-2012 at 09:21

Quote: Originally posted by Pulverulescent

If one is unable, or unwilling (too lazy) to prepare strong HNO3 by distillation then nitration by H2SO4/nitrate-salt mixtures would seem to be a doable alternative, poor though it is . . .

Use of fuming nitric acid for nitrations is unprofitable in a way as if one can perform successful nitrations using nitrate salts then why not?The extremely precious fuming counterpart can be used for other syntheses...
I wonder the difference in yield using nitrate-salt mix. and using aezotropic nitric acid method. Anyone?

Pulverulescent - 12-3-2012 at 09:39

Quote:

The extremely precious fuming counterpart can be used for other syntheses...

If you have the glass, the alkali nitrate and the sulphuric acid why would it be precious?

Bhaskar - 12-3-2012 at 09:42

Good point, my bad.
I was talking about it's price for buying as a reagent.

garage chemist - 12-3-2012 at 11:14

There are several reasons why nitrating mixtures with "real" HNO3 are preferable over those from sulfuric acid and an alkali metal nitrate, even despite the increased water content when using 70% HNO3.

One is recyclability of the spent nitrating acid, a decisive factor in industry. If you use a nitrate salt, the spent acid will be full of alkali metal bisulfate and cannot simply be boiled down to
recycle the sulfuric acid. With HNO3, the spent acid can be freed of HNO3 and water by distillation and used over and over again, without losses except for what adheres to your filtered solid product, or almost no losses in the case of liquid nitrated products.

Another one is viscosity of the nitrating mix. A solution of KNO3 in H2SO4 is a relatively thick and slurry-like, and does not so easily wet things like cellulose, forming a mass that has to be kneaded. With HNO3, the nitrating acid is thin and fluid, rapidly and easily wetting all of the precursor.
Have you ever directly compared nitrocellulose made from acid and nitrate mix with the product from conc. H2SO4 and 65% HNO3? The product made with HNO3 is significantly better, despite more water being present in the nitrating mix!

phlogiston - 12-3-2012 at 12:28

Would adding a little water to the mix to decrease the viscosity help, then?

Bot0nist - 12-3-2012 at 12:30

Using ammonium nitrate instead of potassium greatly reduces the viscosity, IME.

quicksilver - 12-3-2012 at 12:51

Just from a discussion standpoint; I see every reason (if you do have the glass) to distill your acid. One maintains a variety of controls from that agenda.

Sometime back there was a discussion as to the efficiency of varying alkali nitrates used in a mixed acid distillation. Of course there is a small learning curve also as the mixture of a solid nitrate and another acid such as sulfuric acid appears to provide a better yield when the distillation is started from a clear solution of both materials (beginning from a weighted solid simply dumped into a flask of sulfuric creates several problems). I did a small experiment a long time back testing the proportionate yield of nitric acid with a vacuum applied, cooling in the receiving flask and cooling in the condenser (as well as condenser design) and removed.
However, given that a fairly equipped lab would have more than kitchen utensils available, the resultant nitric acid eventually could be substantially inexpensive with a marginal initial investment.

caterpillar - 12-3-2012 at 14:38

I used satls and sulphuric acide where it was possible. Moreover, salts plus aforementioned sulphuric acid are more powerful nitrating agents than HNO3 + H2SO4. The only reason to prepare conc HNO3 by distillation is syntesis of such compounds like RDX and HMX, that are destoyed by H2SO4.

Pulverulescent - 12-3-2012 at 15:28

Quote:

The only reason to prepare conc HNO3 by distillation is syntesis of such compounds like RDX and HMX, that are destoyed by H2SO4.

Mixed acid nitration gives higher yields and better purity . . .
Only one more reason!

AndersHoveland - 12-3-2012 at 17:25

The reason that 70% HNO3 is typically used is because it is completely soluble. Potassium nitrate can have some trouble dissolving in concentrated sulfuric acid. The potassium bisulfate that forms apparently does not have a very high solubility under these conditions because of the common ion effect.

Anhydrous (100 percent concentrated) nitric acid would be the most ideal, but usually these high concentrations of nitric acid are not readily obtainable. Higher concentrations of HNO3 are more difficult to make, and have less stability, having a tendancy to irreversibly decompose in the presence of light. Anhydrous lithium nitrate would also be a good choice, but this is more expensive and not commonly available.

Ammonium nitrate is sometimes used in nitrations because it has such a high solubility and readily dissolves, but in the presence of concentrated sulfuric acid, some of the NH4NO3 can be dehydrated and decompose into nitrogen oxides, reducing yields.

Bhaskar - 12-3-2012 at 21:58

Quote: Originally posted by phlogiston

Would adding a little water to the mix to decrease the viscosity help, then?

Depends on what salt you are using, but it will lead to dilution of the nitration mix. and destroy it. That's what we are talking about, the use of anhydrous nitric acid is preferable to it's aezotropic counterpart as it contains less water.

Bhaskar - 12-3-2012 at 22:03

Won't the salt-acid mix. reduce the purity of the nitrated product and can pose problems for it's synthesis?

AndersHoveland - 13-3-2012 at 00:27

Quote: Originally posted by Bhaskar

Won't the salt-acid mix. reduce the purity of the nitrated product and can pose problems for it's synthesis?

Potentially yes, but typically this is not so much of a problem. It depends on what the nitrated product is. If making cellulose nitrate, the salt can be rinsed out when it is neutralised with sodium bicarbonate solution. If making nitroglyerin, the nitroglycerin separates out as oily droplets which the salt is not soluble in.

caterpillar - 13-3-2012 at 00:28

Quote: Originally posted by Pulverulescent

Quote:

The only reason to prepare conc HNO3 by distillation is syntesis of such compounds like RDX and HMX, that are destoyed by H2SO4.

Mixed acid nitration gives higher yields and better purity . . .
Only one more reason!

It is a strange idea. Who told you it? How do you determine this yield? It can be calculated in different ways. You need generally more H2SO4, and according to it yield will be, yeah, lower. But mixture of salt + H2SO4 is more powerful nitrating agent- less amount of water, and this factor usually increase yield. All sulphates can be deleted by water- they are not reason for impurity.

Bhaskar - 13-3-2012 at 01:55

Quote: Originally posted by AndersHoveland

Quote: Originally posted by Bhaskar

Won't the salt-acid mix. reduce the purity of the nitrated product and can pose problems for it's synthesis?

That's a relief to hear. Nitroglycerine isn't soluble in water right?

AndersHoveland - 13-3-2012 at 02:09

Quote: Originally posted by Bhaskar

Nitroglycerine isn't soluble in water right?

That is correct, it is not.
But nitroglycerin can very slowly hydrolyse in water, especially if there is plenty or residual acid left in the aqueous solution. This is a charactaristic of all esters (nitroglycerin is a nitrate ester of glycerol).

The solubility of nitroglycerin in ethanol is reportedly 1 part nitroglycerin to 3.2 parts by weight ethanol.

Bhaskar - 13-3-2012 at 05:39

Quote: Originally posted by AndersHoveland

Quote: Originally posted by Bhaskar

Nitroglycerine isn't soluble in water right?

So after the nitration of glycerol, I am supposed to wash it in water to clean it, right?
I want to make my first nitroglycerine so that's why I asked...

quicksilver - 13-3-2012 at 12:05

Wash with distilled saline water approximate 2% weight stabilizer,, 5% anti-acid (Naoum, Nitroglycerin & Nitroglycerin Explosives). The allow exposure for approx 48-96 hours, replace, until the NG is neutral to litmus. Upon completion, store in absorbent with anti-acid and stabilizer (Diphenylamine, urea,centrelite, etc, etc) or, if to be kept as liquid, retain under saline water with anti-acid (sodium carbonate, etc).

NOTE: Neutral means 7 - not 7.5. The "salt water" will clean up the NG to a very clear material, pulling any water from it.Stabilizers should not be applied too heavily. There is even a B vitamin that will act as a stabilizer (Betaine).

[Edited on 13-3-2012 by quicksilver]

Bhaskar - 13-3-2012 at 20:00

Quote: Originally posted by quicksilver

Wash with distilled saline water approximate 2% weight stabilizer,, 5% anti-acid (Naoum, Nitroglycerin & Nitroglycerin Explosives). The allow exposure for approx 48-96 hours, replace, until the NG is neutral to litmus. Upon completion, store in absorbent with anti-acid and stabilizer (Diphenylamine, urea,centrelite, etc, etc) or, if to be kept as liquid, retain under saline water with anti-acid (sodium carbonate, etc).

NOTE: Neutral means 7 - not 7.5. The "salt water" will clean up the NG to a very clear material, pulling any water from it.Stabilizers should not be applied too heavily. There is even a B vitamin that will act as a stabilizer (Betaine).

[Edited on 13-3-2012 by quicksilver]

thank you!
Are stabilizers compulsary even for short term purposes?

Bot0nist - 14-3-2012 at 06:16

It is not a bad idea, but is really important in longer term storage, IIRC. We are getting way off topic here though. I'm not scolding anyone, because I know more than anyone how easy a discussion can stray. For future searcher's sake can we please post NG questions in one of the many corresponding threads.

quicksilver - 14-3-2012 at 09:07

Quote: Originally posted by Bot0nist

It is not a bad idea, but is really important in longer term storage, IIRC. We are getting way off topic here though. I'm not scolding anyone, because I know more than anyone how easy a discussion can stray. For future searcher's sake can we please post NG questions in one of the many corresponding threads.

Personally, I agree. - with one caveat. Some materials are VERY tough to get the acid out and neutralized. These are OFTEN crystalline shaped materials, but not always. IMO - the cardinal rule is to get that synthesis to a 7 (neutral) pH. If the material is (especially) a crystal or a liquid - - I would put a stabilizer in there IF (1) I was making more than I was going to use that moment, (2) I would not consider making a large amount - period. (3) - one of the most effective methods of getting acid out of a crystal is re-crystallization.

NG (again IMO) should be made in very small amounts. It's an easy material to make - but it can also go badly very quickly from all sorts of things - even UV light.
There are few experiments that you can't do with a few milligrams other than end up blind or learn to type with one hand.
I am not a young guy. Frankly I am a Hell of a lot older than the overwhelming majority of people on this Forum. I also know that I am not the wisest. But I do know that all ten fingers are typing this now. The interest in Energetic Materials is similar to an interest in firearms. The first thing and the last thing to learn and continue with - is how not to fuck up while having a hobby that has and will continue to change and DEMANDS maturity or extracts a vicious price for a blunder.

TheMessenger - 17-3-2012 at 19:55

Quote: Originally posted by niertap

In most energetic material synthesis' 70% HNO3 is used. Why doesn't everyone just use KNO3 or NaNO3? There will be a considerable amount of water that will not have to be removed from the reaction, decreasing the amount of H2SO4 needed.

A 3.8L jug of 98% H2SO4 can be purchased from the hardware store for around $25. Why would someone not in an analytical setting use anything different?

All of us who distill our own HNO3 use 90%+ concentration. Reactions done with HNO3 instead of nitrate salts have 200-300% high yields. ETN synthesized with fertilizer grade KNO3/H2SO4 results in a yield of around 1E > 0.5ETN while ETN synthesized with HNO3/H2SO4 results in a yield of 1E > 2ETN by weight.

You have been SERVED by TheMessenger.

BromicAcid - 17-3-2012 at 20:27

Quote: Originally posted by caterpillar

How do you determine this yield? It can be calculated in different ways. You need generally more H2SO4, and according to it yield will be, yeah, lower.

Can't speak for Pulverulescent but I can't think of any reasonable chemist that would calculate yield based on anything BUT the key raw. In which case concentration be damned, the amount of sulfuric acid would not make a lick of difference unless you can't recover your product out of it in which case it really wouldn't count as yield would it?

I certainly agree on the front of purity however as a star in favor of using nitric acid. In terms of energetics as far as I know you always want to maintain the best purity you can in order to maintain consistency of your final product and reduce the chance of decomposition (unintentional decomposition of course). Certainly most of us have heard of a few percent of this or that stabilizing or conversely catalyzing the decomposition of a compound. Be your cation ammonium, sodium, potassium, etc. that unquestionably introduces another variable into the equation.

Edit: Counterpoint to the title, Dinitrogen pentoxide is an excellent nitronium ion source, why not use it?

[Edited on 3/18/2012 by BromicAcid]

AndersHoveland - 17-3-2012 at 22:06

Quote: Originally posted by BromicAcid

Counterpoint to the title, Dinitrogen pentoxide is an excellent nitronium ion source, why not use it?

Yes, N2O5 is an excellent nitronium cation source, ideal in several ways. But it also has some disadvantages: it is a relatively sensitive explosive, it is unstable and will decompose after about 2 days at room temperature, it is difficult to prepare and is not commercially available.

For these reasons, nitronium tetrafluoroborate, NO2BF4, is usually preferred. Another less commonly thought of disadvantage to N2O5 relative to NO2BF4 is that the nitrate anion itself can actually be problematic in certain extreme reactions.

Quote: Originally posted by AndersHoveland

NO2BF4 must be used, not N2O5, because the inermediate appears to be reactive toward the nitrate anion, which then results only in oxidizing the amino group to a nitro group, rather than formation of a new dioxytetrazine ring.
http://www.sciencemadness.org/talk/files.php?pid=233466&...

For using N2O5, one possible method involves using pressurised liquid CO2 as the solvent.
https://www.sciencemadness.org/whisper/viewthread.php?tid=17...
http://www.sciencemadness.org/talk/viewthread.php?tid=17304
The danger, of course, would be that as soon as the mixture was depressurized, the CO2 would immediately boil out, leaving the solid explosive N2O5. The sudden decompression could potentially trigger detonation.
SO2 would be much easier to liquify, but N2O5 would probably oxidize the SO2 to nitrosyl pyrosulfate, (NO)2S2O7

Preparation of Nitronium Tetrafluoroborate
Nitronium tetrafluoroborate is a useful nitrating agent which has several advantages over nitric acid. There are also some nitration reactions where the presence of the nitrate ion would prevent the desired product from forming. Nitronium tetrafluoroborate is a solid ionic compound, with the formula NO2(+) BF4(-). Nitronium tetrafluoroborate can be prepared by adding a mixture of anhydrous hydrogen fluoride gas and boron trifluoride to a solution of either highly concentrated nitric acid or nitrogen pentoxide dissolved in nitromethane.

WARNING: Rubber gloves, an apron, and a plastic face mask are strongly recommended. All operations should be carried out in a hood. If hydrogen fluoride comes in contact with the skin, the contacted area should be thoroughly washed with water and then immersed in ice water while the patient is taken to rushed to the emergency department. Burns caused by hydrogen fluoride may not be noticed for several hours, by which time serious tissue damage may have occurred.
Note: Any operations involving liquid hydrogen fluoride must be carried out with equipment resisting hydrogen fluoride, such as fused silica, polyolefin, monel steel, or teflon. Kel-F grease is recommended for ground-glass joints. Nitronium tetrafluoroborate slowly attacks silicone stopcock grease, causing air to enter the flask. After completion of the reaction, all equipment should be washed with plenty of water.

A 1-Liter three-necked polyolefin flask is provided with a short inlet tube for nitrogen, a long inlet tube for gaseous boron trifluoride, a drying tube, and a magnetic stirring bar. The flask is immersed in an ice-salt bath and flushed with dry nitrogen. Under a gentle stream of nitrogen and with stirring, the flask is charged with 400 ml. of methylene chloride, 41 ml. (65.5 g., 1.00 mole) of red fuming nitric acid (95%), and 22 ml. (22 g., 1.10 moles) of cold, liquid, anhydrous hydrogen fluoride. 5. It is convenient to condense anhydrous hydrogen fluoride, b.p. 19.5°, from a cylinder into a small calibrated polyolefin flask immersed in a mixture of dry ice and acetone. As hydrogen fluoride is very hygroscopic, it should be carefully protected from atmospheric moisture, preferably by maintaining an atmosphere of dry nitrogen over it, otherwise by means of a drying tube. The hydrogen fluoride is then simply poured into the reaction flask.

Gaseous boron trifluoride (136 g., 2.00 moles) from a cylinder mounted on a scale is bubbled into the stirred, cooled reaction mixture. (The temperature of the reaction is not critical, but the reaction is slower at higher temperatures because of the lower solubility of boron trifluoride in the solvent). The first mole is passed in rather quickly (in about 10 minutes). When approximately 1 mole has been absorbed, copious white fumes begin to appear at the exit, and the rate of flow is diminished so that it takes about 1 hour to pass in the second mole; even at this slow rate, there is considerable fuming at the exit. After all the boron trifluoride has been introduced, the mixture is allowed to stand in the cooling bath under a slow stream of nitrogen for 1.5 hours. The mixture is swirled, and the suspended product is separated from the supernatant liquid by means of a medium-porosity, sintered-glass Buchner funnel. Note that since free hydrogen fluoride is no longer present, filtration can be carried out with glass or porcelain equipment.

The gooey solid remaining in the flask is transferred to the funnel with the aid of two 50-ml. portions of nitromethane. The solid on the funnel, nitronium tetrafluoroborate, is washed successively with two 100-ml. portions of nitromethane and two 100-ml. portions of methylene chloride. In order to protect the salt from atmospheric moisture during the washing procedure, suction is always stopped while the salt is still moist. The moist salt is transferred to a round-bottomed flask and dried by evaporating the solvent. At the end of the procedure the flask can be gently heated to 40–50°C (Nitronium tetrafluoroborate is thermally stable up to 170°. Above this temperature it starts to dissociate into nitryl fluoride and boron trifluoride.) The yield of colorless nitronium tetrafluoroborate is 85–106 g. (64–80%) It is stored in a wide-mouthed polyolefin bottle with a screw cap. The edge of the cap is sealed with paraffin wax after it is screwed on. Nitronium tetrafluoroborate is very hygroscopic. It is stable as long as it is anhydrous, but it is decomposed by moisture, and all transfers should be in a dry box.

Nitronium tetrafluoroborate slowly attacks polyethylene and polypropylene, but apparatus made of these materials will last for several preparations of the salt.

The last part of the procedure can be used to purify nitronium tetrafluoroborate that has picked up water on standing. The impure salt is washed twice with nitromethane, twice with methylene chloride, and is dried under reduced pressure.

Making Boron Trifluoride BF3
Boron trifluoride is a very toxic gas, which readily reacts with water to form metaboric acid and Fluoroboric acid, which then can further hydrolyze if excess water is present. It is a strong fluoride ion abductor, meaning it will pull a fluorine atom from many covalent compounds to form the tetrafluoroborate anion (BF4-), while leaving a positively charged cation. However, boron trifluoride is not as powerful of an abductor as antimony pentafluoride, as demonstrated by its unreactivity towards trichlorofluoromethane.

Boron trifluoride may be prepared by heating a mixture of boric oxide and calcium fluoride with concentrated sulfuric acid. It may also be prepared by mixing 5 parts of potassium borofluoride (KBF4) with 1 part finely powdered boric oxide, then heating with concentrated sulfuric acid. The boron trifluoride, can be collected over mercury. Potassium borofluoride may be produced by heating together 2 parts boric acid, 5 parts CaF2, and 10 parts conc H2SO4. The liquid is then cooled and filtered, and a solution of a potassium salt is added. Potassium borofluoride precipitates out, and may then be recrystallized from hot water. By this method it can be prepared as anhydrous hexagonal crystals. If it is prepared instead from hydrofluoric acid, boric acid, and K2CO3, a gelatinous mass forms instead, which however forms cubic octahedral and cubic dodecahedral crystals when heated to 100degC. Boron trifluoride also results from heating a mixture of boric oxide and calcium fluoride to a white heat in an iron pipe. Heating solid borofluorides to red heat, boron trifluoride is evolved, leaving behind metal fluorides.

[Edited on 18-3-2012 by AndersHoveland]

AndersHoveland - 24-8-2013 at 08:14

Here is a different method for preparation of nitronium tetrafluoroborate:

Quote:

Prepare a mixture of 8.8 g of hydrofluoric acid and 40 g of ethylnitrate in 300 mL of dry nitromethane. Place this mixture in a salt-ice bath and cool to 10 C. Slowly add 100 g of boron trifluoride to the mixture while keeping the temperature between 10 and 15 C. A few minutes after the addition filter to collect the crystals of nitronium tetrafluorate that form. Wash the crystals with two 400 mL portions of a 1:1 mixture of methylene chloride and nitromethane, and then with one 400 mL portion of methylene chloride.

The ethyl nitrate being made by nitration on ethanol, of course.
Not sure what exactly they mean by "hydrofluoric acid", might be better to pass the equivalent amount of HF gas into the reaction (using plastic containers and tubing, and taking the appropriate precautions).

[Edited on 24-8-2013 by AndersHoveland]