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Author: Subject: Nitric Acid Synthesis
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[*] posted on 8-7-2008 at 00:26


I tried a similar thing, I mixed roughly 55 ml of sulphuric acid with 55 grams of NaNO3, which means quite a large excess of H2SO4. From this I obtained 23 ml of densely fuming nitric acid with a yellow color and a density of appr. 1.5 g/ml. Total yield (taking into account the amount of NaNO3) is between 75% and 80%.

I noticed that I had a steady production of nitric acid (1 drop per second) with a temperature around 83 C. At a certain point, the temperature was rising, and the mess in the heated flask started foaming much more. At a certain point, the temperature rose to 90 C. At that point I stopped, realizing that I may not get all HNO3, but also leaving the higher boiling azeotrope in the flask.

So, I think that you should not try to squeeze out the single last drop of nitric acid from the mix. Only the part, which boils well below 90 C is the concentrated stuff, at higher temperature, more water boils over. The water is from the sulphuric acid. Even the concentrated stuff frequently is only 92% or so, especially if technical grade material is used.

The second is that using an excess amount of sulphuric acid makes for a much easier procedure, keeping more of the water behind in the hygroscopic mix.

You could of course change receiver flasks when the temperature rises above 90 C. The last fraction then is more dilute acid, which may be useful as well. I did not bother about that final fraction. I already have a few liters of lab grade 65% acid, I only was interested in the pure acid.

The chems I used were fertilizer grade NaNO3 and drain cleaner (for the dutch members: Mega ontstopper), which is said to contain 97% H2SO4, and has a very light brown color. So, I did not use anything fancy.

I also noticed that initially, a lot of NO2 is formed, but after a while the produced acid is much cleaner. I think that this initial formation of NO2 is due to impurity in the fertilizer. The color of the granules of NaNO3 is off-white, some granules even are brown/yellow. I think this is due to organic impurities, which are oxidized when the mixed is heated, and once all this crap is oxidized, the cleaner acid is distilled. But even then, I never obtained colorless acid, the drops remain golden yellow.

So, for me, I am satisfied by now. For each 50 ml of drain cleaner I can make 20 ml of yellow fuming nitric acid of density 1.5. NaNO3 is not an issue at all for me, I purchased a 25 kg bag for EUR 10, which can be purchased at any bigger shop for gardening (for dutch members: Welkoop). So, the cost of the nitric acid is determined by the cost of the sulphuric acid I use (appr. EUR 15 per liter). Total cost for me is EUR 0.75 per 20 ml.




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[*] posted on 8-7-2008 at 02:50


I'm also planning to make some nitric acid >90%. For it's storage a DURAN 50mL bottle with a PTFE lined, red cap will be bought.

But the danger of the procedure kind off holds me from doing it. I start from nitric acid and sulfuric acid. My question is, how much NO2 evolves? I can only do this inside, and I can lead the fumes through NaOH-solution (does this absord all NO2?) through the vaccuum adapter, but when opening the apparatus, I would have a bit of a problem. I can ventilate the garage fairly well, but not like outside or a fume hood.
My second question is, how dangerous is highly concentrated nitric acid? For example, I have a wooden bench (somewhat hard wood, coated with some plastic), so what if the flask breaks? Fire?
Finally, can the acid be simply discolored with urea? I thought so, and I think there's only evolutio of N2, CO and water.
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[*] posted on 8-7-2008 at 03:28


Very little NO2 should be produced so long as you heat it gently. When mixing the nitric and sulphuric acids remember that there is a production of heat from the dilution of the conc. sulphuric acid. Supposedly the yellow/red acid can be discoloured by pulling a vacuum (if using a water aspirator remember a drying tube between the flask and the aspirator), or by bubbling through dry air.

Bubbling the dry air dissolves any NO2 in the remaining water producing HNO3 (by providing the oxygen for this reaction: 4NO2 + O2 + 2H2O => 4HNO3). Therefore I would suggges this method to be better as you are not only getting rid of the NO2, but you are also increasing the concentration of your acid :P
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[*] posted on 8-7-2008 at 04:55


Sadly, the oxidation of NO2 with oxygen does not work for very high concentration of the acid. But bubbling through dry air (free of dust!!!) does help making a pure acid. It drives off NO2, dissolved in the acid. But be prepared, this is a very slow process.

Almost 100% nitric acid is very dangerous stuff. Concentrated (96%) sulphuric acid is a toy for little kids, when compared with pure nitric acid. The acid I have is fuming intensely. When I open the bottle, then an amazingly thick white cloud is produced and pouring it out of the bottle also produces LOTs of fume. Just for fun, I dropped some of the acid on a piece of paper, and this results in a brown cloud of NO2, and a hole in the paper, but no fire was produced. I can imagine though that at larger quantities this could lead to fire.

I did the synth inside, having a bucket of water nearby, just in case something went wrong. I also limit the quantities to 50 grams or so, and use a 500 ml RBF, due to the excessive foaming during the procedure.

Some months ago, I also did the experiment, as described by Jor, using 65% nitric acid and excess concentrated sulphuric acid. This method produces a somewhat cleaner acid than the method with NaNO3, but now I have both acids and can compare them, I have the impression that the stuff, made from NaNO3 is more concentrated. The stuff I obtained from the 65% acid does fume in air, but the stuff I obtained from the NaNO3 is fuming insanely. It really is remarkable.

I also noticed another, scary thing. My little bottle, with nitric acid, which I made a few months ago had a large pressure inside. This is a thick-walled 15 ml bottle, filled with well over 10 ml of acid, and in the little amount of air above this, there was a lot of pressure. Does highly concentrated HNO3 decompose, giving oxygen? If this is the case, then having this stuff around is exceptionally dangerous, unless you release the pressure every few weeks, or unless you leave quite some NO2 in the acid, which may get into equilibrium with the oxygen and some water.

Any expert views on this issue of pressure buildup? I only have my own experience, but I would like to have more information from others. Are there people out there with similar experiences?




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[*] posted on 8-7-2008 at 05:28


what are you usign to heat the boiling flask? open flame or an oil bath.. or is it another heat source?

about leaving the nitir cacid in the bottle I think that will not suppose a problem because the elevated pressure will reduce the fuming and thus equalibrium will be rached between the fuming and the rising pressure.
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[*] posted on 8-7-2008 at 05:34


I use a heating mantle for heating the flask. This gives a more evenly spread heat, and I can adjust the heating with that.

Your answer about the pressure buildup is not a real answer to my question. I have the strong impression that the pressure buildup is not simply because of evaporating of the acid, but because of decomposition of the acid. I have read somewhere that highly concentrated nitric acid can decompose to NO2, O2 and H2O. The O2 is what scares me. This leads to excessive pressure buildup. The other produces simply dissolve in the acid.




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[*] posted on 8-7-2008 at 06:04


Dont you think that the increased pressure will stop the decomposition of the acid?

and what if i dont have a heating mantle?! what would you suggest to use?
Im kind of afraid to use an oil bath. you know .. nitric acid + oil = bad news.
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[*] posted on 8-7-2008 at 06:07


Distilling the 100% acid is supposed to form some water and N2O5, which breaks down into (N2O4<=>NO2) and oxygen, so it does seem reasonable that on storage oxygen would form. Keeping the acid cool, near 0 C, would slow both its decomposition and that of the N2O5, which would react with the water formed to give HNO3 again.
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[*] posted on 8-7-2008 at 07:01
WFNA Storage


I found an interesting article titled, "Investigation of Effects of Additives on Storage Properties of Fuming Nitric Acids", at this link: http://naca.central.cranfield.ac.uk/reports/1952/naca-rm-e52... WFNA pressures are a magnitude greater than RFNA. I also saw another reference that said to open bottles to release pressure. (www.gcal.ac.uk/sls/Bio/healthandsafety/sa5.html). The storage problem occurs because the liberated O2 is not soluble in the acid.
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[*] posted on 8-7-2008 at 10:30


Wow, this stuff is extremely corrosive I think! :

http://youtube.com/watch?v=zWK8pl0MmGM&feature=related
http://www.youtube.com/watch?v=Y5HSO5HYnG4&NR=1
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[*] posted on 8-7-2008 at 11:24


Yes, it is VERY corrosive. I did another test. I took 0.2 ml of the acid (just a few drops in the bottom of a test tube) and dripped in some isopropyl alcohol. The effect is stunning. Strong crackling noise, near-explosion :o. This is totally different from mixing alcohols with stuff like 53% HNO3.

I also did a test with NaCl, which is quite inert. But even HCl reacts vigorously with highly concentrated HNO3. It reacts with heavy foaming and an orange/yellow gas mix is produced. It also reacts violently with water. A drop of water, added to the acid, results in boiling.




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[*] posted on 8-7-2008 at 11:41


Indeed concentrated HNO3 is nasty stuff I've gotten small amounts of nitration mix from 70% nitric and 98% sulfuric in 50/50 ratio begin to turn my gloves yellow and weak.

Surprisingly it takes upwards of a minute for my 70% nitric acid to have an effect on wood .

It would appear as if the corrosive power of nitric acid doesn't increase in equal increments to corresponding concentration but grows almost exponentially or at least between azetropic concentration and 100%.




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[*] posted on 8-7-2008 at 12:02


@ Klute, the hydrogen peroxide trick is something I do on a routine basis to stretch my nitric acid consumption when I dissolve up things in aqua regia. It prevents excessive release of nitrogen oxides and reforms HNO3 in situ. I've never had an explosion or even intense reaction using it, ever so I hear about your experience and wonder if there was some other factor involved. I do not have sulfuric present, ever. Only 50% HNO3 and 30% HCl in my aqua regia. I only add a few milliliters at a time.



@ Woelen, the best bottle for white or red fuming nitric acid is one made of ETFE with a Tefzel cap. They are pricey for new ones, about 300USD for a 2L but well worth it! They'll store oleums, 70% HF, any strength nitric, stock dichromate cleaning solutions and many other corrosives. My friend uses those bottles also as pressure reactors. They are good to about 4 bar as I recall, hopefully he'll give his impression on it. The difference between even azeotropic HNO3 and red fuming nitric is immense. Azeotropic nitric does not hurt when it gets on my hand unless I leave it there over long. The 90%+ acid hurts instantly--once I was pipetting some out of said 'teflon' bottle when I moved the pipette over my exposed, ungloved wrist. One drop fell and landed squarly in the middle of my exposed palmup wrist. It felt as if a cigarette were put out on it. I still have a small scar there from that foolish mistake.




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[*] posted on 8-7-2008 at 12:05


This also can be explained. Azeotropic or more dilute HNO3 is largely split in ionic species NO3(-) and H3O(+). NO3(-) is MUCH more stable than HNO3. Pure HNO3 is mostly molecular and only contains a small amount of ionic species, and the species present in pure HNO3 are NO3(-), H3O(+), NO2(+). The latter is an exceptionally reactive species, even more so than molecular HNO3.

This increased reactivity is well known for more strong acids. E.g. 60% HClO4 is surprisingly inert, not even boiling acid is capable of oxidizing iodide to iodine at an appreciable rate. The azeotropic mix with 72% HClO4 still is quite inert and is an excellent non-coordinating acid. But at higher concentrations, the acid rapidly becomes more reactive, and 100% HClO4 explodes on contact with paper, wood, etc.
A similar effect to a lesser extent exists for pure H5IO6 and its solutions in water.

Many anionic species are resonance stabilized, e.g. in NO3(-) all oxygens are equivalent and the single charge is distributed over the entire ion. In an acid like HNO3 this symmetry is broken, one of the oxygens carries the hydrogen atom, and the two others have double bonds to the nitrogen atom. This latter structure is MUCH more reactive than the symmetric nitrate ion.




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[*] posted on 8-7-2008 at 13:03


It is indeed very corosive, I wonce got a few drops splattered over my arm, I didnt make it to the bathroom to wash it off, ended up with 4 really deep black holes in my arm that took several weeks to heal.
very nasty stuff.

PS. can anybody answer my question about the heating?
how would you recomend heating the flask (a mantle is not an option)
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[*] posted on 8-7-2008 at 13:20


Quote:

PS. can anybody answer my question about the heating?
how would you recomend heating the flask (a mantle is not an option)


What kind of heating source do you have? If you don't have a heating mantle I would suggest a laboratory hotplate or if not just a standard electric hotplate.

Then use a sand bath, make sure to use clean washed sand. Maybe run it through a screen a few times wash it and stick it in the oven at 350 degrees for a few hours. Then pun it in a suitable container on top of the hotplate and put the boiling flash about 1/3 to 1/2 the way in I doubt sulfuric or nitric acid will react with that.


I would still wear protection if in some freak accident nitric acid gets on hot sand or anything hot for that matter there will be a lot of Nox fumes.




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[*] posted on 8-7-2008 at 17:11


Quote:
Originally posted by Jor
I'm also planning to make some nitric acid >90%. For it's storage a DURAN 50mL bottle with a PTFE lined, red cap will be bought.


Be careful which plastic it is stored with. It will react with many plastics. The plastic lids in containers used for concentrated nitric acid can get munched right through, like this at the bottom (kind of a silly page, but the image demonstrates the point). But PTFE is unreactive toward the nastiest chemicals, so it should work.

Quote:
But the danger of the procedure kind off holds me from doing it. I start from nitric acid and sulfuric acid. My question is, how much NO2 evolves? I can only do this inside, and I can lead the fumes through NaOH-solution (does this absord all NO2?) through the vaccuum adapter, but when opening the apparatus, I would have a bit of a problem. I can ventilate the garage fairly well, but not like outside or a fume hood.


I wouldn't do this inside. You can also neutralize NO2 vapors with a cheap NH3 solution.

Quote:
My second question is, how dangerous is highly concentrated nitric acid? For example, I have a wooden bench (somewhat hard wood, coated with some plastic), so what if the flask breaks? Fire?


Highly concentrated HNO3 is dangerous. It will nitrate organic materials on its own to produce nitrate esters and nitro compounds, like with wood. If the nitric acid contacts the wood, it is a definite fire risk. Nitrate esters are prone to spontaneous decomposition, ignition and explosion in the presence of acidity.

Quote:
Finally, can the acid be simply discolored with urea? I thought so, and I think there's only evolutio of N2, CO and water.


Urea will react with both HNO2 and nitrogen oxides. Urea will react with HNO2 in mole ratio of 2:1 at regular temperature to form NH4NO2 and HNCO. Heating this forms: (NH4)2CO3, N2, and CO2. With an excess of HNO2, only N2 and CO2 result: (NH4)2CO + 2 HNO2 -> CO2 + 2 N2 + 3 H2O. Urea will also decompose nitrogen oxides led into its solutions optimally at 80 deg. at best in 40% conc. (Gmelin C [D1] 419).

I have no specific method for using urea with highly concnd. HNO3. With conc. HNO3 urea yields explosive barely soluble urea nitrate. From Beilstein 1, 1291: highly conc. HNO3 decomposes urea nitrate as follows: CO(NH2)2.HNO3 + HNO3 = CO2 + N2O + NH4NO3 + H2O.

I've added some urea to fuming HNO3 (1.52), and it did clear all of the reddish gas and yellow color in a fizzing reaction leaving the acid afterwards basically colorless and fuming only white, though a small portion of the urea looked unreacted, but this dissolved after gentle heating. Not sure about contamination with this method though.

For ridding of NO2 gases, Gmelin N [3-4] p. 961 says: (preferably warm) air or CO2-stream is lead into the acid, this will lead the gases away. To keep highly concd. acid completley colorless, distill in a glass apparatus at 45 deg. under reduced pressure of 15 mm Hg. They say the most effective method for removing last traces of NO2 is from leading in an ozonized-O2 stream and then distilling in a vaccum.

Quote:
Originally posted by woelen I also noticed another, scary thing. My little bottle, with nitric acid, which I made a few months ago had a large pressure inside. This is a thick-walled 15 ml bottle, filled with well over 10 ml of acid, and in the little amount of air above this, there was a lot of pressure. Does highly concentrated HNO3 decompose, giving oxygen? If this is the case, then having this stuff around is exceptionally dangerous, unless you release the pressure every few weeks, or unless you leave quite some NO2 in the acid, which may get into equilibrium with the oxygen and some water.

Any expert views on this issue of pressure buildup? I only have my own experience, but I would like to have more information from others. Are there people out there with similar experiences?


It is known that HNO3 decomposes from thermal as well as UV-radiation and light, especially in combination: 4 HNO3 --> 4 NO2 + 2 H2O + O2. This is also why nitric acid is typically stored in dark bottles. It is also possible to end up with Cl2-contaminated acid from that one obtained from H2SO4 and alkali nitrates containing chloride contaminants. Other than those possibilities, I'm not sure why there should be a pressure I used to have a red fuming nitric acid in small brown glass bottles stored in the dark with a volume of the acid similar to the bottle stored for years with no bursting, violent pressure accumulation.
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[*] posted on 8-7-2008 at 23:50


I have been roeading around a little on this subject, and I found that red fuming nitric acid is much more safe on storage than white fuming nitric acid. Red fuming nitric acid hardly suffers from pressure buildup, while white fuming nitric acid does have this problem. The red fuming nitric acid contains a lot of free NO2 and this drives the equilibrium reaction back towards nitric acid. The decomposition of white fuming nitric acid appafrently cannot be halted, nor prevented. Cool storage at 0 C makes decomposition slower, but it will not stop. When sufficient acid has decomposed and there is sufficient water and NO2 left in the acid, then the decomposition stops, but at that point your acid is not colorless anymore and I would not call it 'white fuming nitric acid' anymore.

Probably, the best is to store the acid as red fuming acid and when it is needed, then a small amount is decontaminated (e.g. by ureu or by bubbling warm dry air through it) just before use.




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[*] posted on 29-7-2008 at 20:45


Because urea nitrate is basically just a double salt of urea and nitric acid, do you think it would be possible to separate the two? If you could that would be an easy way of nitric acid manufacturing.



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[*] posted on 29-7-2008 at 20:47


I would react it with a strong alkali first, to yield the corresponding metal nitrate (these are the most convenient salts for nitric acid synthesis via acidification) and ammonia. But why would you be looking to urea nitrate as a precursor to nitric acid anyway? Usually you'd be using nitric acid to make urea nitrate, not the other way around, unless you can just go buy urea nitrate whenever you want no questions asked.:o

[Edited on 29-7-2008 by kilowatt]




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[*] posted on 29-7-2008 at 21:05


Well, you can make urea nitrate via HCl + NH4NO3. That is one of the only "nitrations" that can be done with that route because of the NH4Cl byproduct. So, if I made urea nitrate with that route, and then separated the HNO3 out, I could start having some real fun. ;)



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[*] posted on 29-7-2008 at 21:12


It would be easier to just use the NH4NO3 to form metal nitrate and use that to make nitric acid. If you are really ambitious you could even use the ammonia that is given off in the first reaction to make more nitric acid.



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[*] posted on 29-7-2008 at 21:15


How would you synth HNO3 from a metal nitrate without sulfuric acid? That is the whole point of doing this multi-step process.



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[*] posted on 29-7-2008 at 21:18


Well I was thinking of doing it with sulfuric acid. I believe urea nitrate could be liable to explode if heated, but I'm really not sure how stable it is. Acidifying NH4NO3 directly to get HNO3 would likely be a bad idea as well.

If you can get HCl I would imagine you can get H2SO4. Have you tried auto parts stores (battery acid) or hardware stores (conc. sulfuric drain cleaner)?

Another idea could be to use copper or lead oxide, hydroxide, or carbonate with NH4NO3 to make copper or lead nitrate. While this reaction is very slow and driven by the loss of ammonia (requires boiling), copper or lead nitrates can be directly decomposed by heating to NO2 and O2 in the correct proportions for making nitric acid if bubbled into water.

[Edited on 29-7-2008 by kilowatt]




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[*] posted on 29-7-2008 at 21:36


I have checked all of the drain cleaners available, but they only contain conc. NaOH. There aren't any auto stores nearby either, or I wouldn't have asked. I got the HCl as muriatic acid from home depot.
Although the copper/lead nitrate sounds promising.

EDIT: I have also been thinking about the electric arc process, but I can't find any neon sign transformers.

[Edited on 7/29/2008 by Zelot]




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