Sciencemadness Discussion Board - Nitric Acid Synthesis

Nitric Acid Synthesis

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woelen - 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.

Jor - 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.

DJF90 - 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

woelen - 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?

hashashan - 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.

woelen - 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.

hashashan - 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.

not_important - 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.

WFNA Storage

bigbigbeaker - 8-7-2008 at 07:01

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.

Jor - 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

woelen - 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

. 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.

crazyboy - 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%.

Fleaker - 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.

woelen - 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.

hashashan - 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)

crazyboy - 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.

Formatik - 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.

woelen - 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.

Zelot - 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.

kilowatt - 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.

[Edited on 29-7-2008 by kilowatt]

Zelot - 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.

kilowatt - 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.

Zelot - 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.

kilowatt - 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]

Zelot - 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]

Picric-A - 30-7-2008 at 01:06

can ammoniium nitrate be used to make nitric acid via H2SO4 + nitrate salt method?
Picric a

OMG - 30-7-2008 at 10:30

Hi everyone, I have been a non member for a while reading and learning chemistry. This is a great place to learn.

I have a question about making nitric from ammonium sulfate.

In a sealed vessel, would it be possible to heat ammonium sulfate to decomposition but keep the temperature below H2SO4's boiling point.. so about 280-290 C. Then have an ozone generator in the reaction chamber to oxidize the ammonia given off? Producing ammonium nitrate which would get converted to nitric acid by the H2SO4.

Essentially the setup would be oxidizing the heck out the ammonium sulfate to end up with nitric and sulfuric acid.

I'm finding it very very difficult to locate any relatively pure quantity of any kind of nitrate. So I'm exploring other methods. I realize that this method would require a lot of ozone to produce any significant amount of nitric acid.

Thanks

crazyboy - 30-7-2008 at 13:17

Quote:

Originally posted by Picric-A
can ammoniium nitrate be used to make nitric acid via H2SO4 + nitrate salt method?
Picric a

Yes but you may have to adjust the amounts.

Quote:

In a sealed vessel, would it be possible to heat ammonium sulfate to decomposition but keep the temperature below H2SO4's boiling point.. so about 280-290 C. Then have an ozone generator in the reaction chamber to oxidize the ammonia given off? Producing ammonium nitrate which would get converted to nitric acid by the H2SO4.

I'm not sure if that method will work at all but if it somehow manges to work you will be getting huge volumes of NOx fumes . It's quite possible all the nitric acid will decompose at 100+ degrees.

kilowatt - 30-7-2008 at 15:16

You'd need one heck of an ozone generator to oxidize enough ammonia to make significant mole quantities of NO2/nitric acid. Ozone generators are normally pretty weeny unless we're talking about a large UV unit. It would be much easier and more energy efficient to use regular air and a catalyst from an old catalytic converter at about 500°C and 50-150psi. The honeycomb catalyst is very easy to cut and make fit into a pipe or other reactor of your making.

ShadowWarrior4444 - 30-7-2008 at 16:00

Quote:

Originally posted by kilowatt
You'd need one heck of an ozone generator to oxidize enough ammonia to make significant mole quantities of NO2/nitric acid. Ozone generators are normally pretty weeny unless we're talking about a large UV unit. It would be much easier and more energy efficient to use regular air and a catalyst from an old catalytic converter at about 500°C and 50-150psi. The honeycomb catalyst is very easy to cut and make fit into a pipe or other reactor of your making.

Nitric Acid can be produced in a useful quantity by using a corona or arc discharge. Corona discharge is employed in many ozone generators and can be easily achieved with a neon sign transformer. (You would use it oxidize the N2 in the atmosphere, not ammonia, as this would waste oxygen.)

A catalytic converter will not work, as it is designed to reduce NOx species to N2 and xO2. (If it is a three-stage converter, as cars employ.)

I do recall that copper can catalytically oxidize ammonia to NOx, however it is so exothermic that the copper will melt. Creative mounting of the copper and flow regulation of the ammonia and oxygen feedstock may be useful.

kilowatt - 30-7-2008 at 17:05

Quote:

A catalytic converter will not work, as it is designed to reduce NOx species to N2 and xO2. (If it is a three-stage converter, as cars employ.)

Automotive catalytic converters use the same catalyst materials (platinum or platinum/palladium) as the Ostwald Process for oxidizing ammonia to NO2. As a demonstration of their oxidizing property just shoot some propane or butane through one of the honeycombs, and it will be oxidized without combustion bringing the catalyst to a bright red heat (has to be hot to start the reaction though). I have done this routinely to check if old catalysts are still active.

ShadowWarrior4444 - 31-7-2008 at 17:57

Quote:

Originally posted by kilowatt

Quote:

A catalytic converter will not work, as it is designed to reduce NOx species to N2 and xO2. (If it is a three-stage converter, as cars employ.)

Of course it oxidizes fuel. Three-way catalytic converters are designed to do three things:

1. Oxidize any unburnt fuel to H2O and CO2.
2. Oxidize CO to CO2
3. Reduce NOx to N2 and xO2

Some older cars may use only a two-way converter, however most if not all modern cars use a three-way. You should check to see which one you have before using it for nitric acid synthesis. Some three-way converters use two chambers as well, one for reduction, one for oxidation.\

Ancillary: Old propane catalytic heaters may be more useful and easier to acquire. (Depending on circumstance.)

[Edited on 7-31-2008 by ShadowWarrior4444]

kilowatt - 31-7-2008 at 18:06

If you use the latter catalyst from a 3 stage converter it should work fine. Every one I have opened has two, presumably one for NOx reduction, and one for the other two functions. I haven't actually done the synthesis yet though, too much else going on.

[Edited on 31-7-2008 by kilowatt]

Contrabasso - 19-8-2008 at 03:52

OK! to stretch the topic a little.

I have access to 60% HNO3 at sensible prices and would like to concentrate some to 95%.

Would a reasonable method be vacuum distilation at a moderate vacuum using an aspirator?

I'm aiming to have a couple of litres of product from a quickfit style apparatus.

ScienceSquirrel - 19-8-2008 at 04:15

Quote:

Originally posted by Contrabasso
OK! to stretch the topic a little.

I have access to 60% HNO3 at sensible prices and would like to concentrate some to 95%.

Would a reasonable method be vacuum distilation at a moderate vacuum using an aspirator?

I'm aiming to have a couple of litres of product from a quickfit style apparatus.

You could try a Google on 'nitric acid distillation', collect up a few ideas and then come back to this board with some idea of what you are trying to do.

woelen - 19-8-2008 at 04:16

Without concentrated sulphuric acid you will never go higher than 68%, due to the azeotrope of water and HNO3. I'm not sure about the exact percentages at other than atmospheric pressure, but I'm quite sure that you will not even come close to 95%.

If you have concentrated sulphuric acid, then you have an easy job. Just mix equal volumes of acid and slowly distill, you do not even need a vacuum distillation, normal pressure is suitable. Keep on distilling, as long as the temperature remains below 90 C. Suddenly, at a certain point, the temperature fairly quickly rises. That is the point where you should stop. You can reuse your sulphuric acid, by boiling it down after the distillation and then again mixing it with the nitric acid.

The acid, obtained in this way is yellow, due to some dissolved NO2. Be aware with the storage of highly concentrated HNO3. I have appr. 30 ml of 90+ % HNO3 and I need to unscrew the cap every few weeks, due to pressure buildup. Not good at all.

ScienceSquirrel - 19-8-2008 at 05:01

There are plenty of recipes on the Internet along the lines outlined by Woelen and they will yield an acid of very high purity.
I can't remember the details now, O level chemistry in the late 1970's, but very concentrated nitric acid is an appreciably more vigorous oxidising agent compared with normal 60 - 70% lab acid.
A small pile of warm, dry sawdust will catch fire on adding a few drops of the concentrated acid.
I think having a look through some old practical books will throw up some interesting demos

grndpndr - 20-8-2008 at 14:29

Quote:

Originally posted by ScienceSquirrel
There are plenty of recipes on the Internet along the lines outlined by Woelen and they will yield an acid of very high purity.
I can't remember the details now, O level chemistry in the late 1970's, but very concentrated nitric acid is an appreciably more vigorous oxidising agent compared with normal 60 - 70% lab acid.
A small pile of warm, dry sawdust will catch fire on adding a few drops of the concentrated acid.
I think having a look through some old practical books will throw up some interesting demos

I am in possesion of at least 50-60yr old Holt-Rinehart-Winston, I believe a college chemistry 100 level lab manual.The glassware etc is rudimentary But the illustrated labs such as the Nitric acid
synthesis using an alchemists type glass retort with a small snug fitting round bottom flask in an ice bath to catch the
distillate.Whats astounding is the list of chemicals assigned to each student.As an amateur you'd be lucky to be able to steal some of these things from a well stocked university lab.
Red white P,uranium ore,pott cyanide?

If I were to start again Id likely use that old chem manual method to low buck small amts 90%+ nitric acid although the same can be done with very crude equipment and patience.

PVC tape ad PVC tubing lasts a few distillations before it gets soft when hot then brittle when cold and needs replaced but it did work for small amounts of fumng red nitric clearly stronger than the 70%.

[Edited on 20-8-2008 by grndpndr]

[Edited on 20-8-2008 by grndpndr]

ordenblitz - 20-8-2008 at 20:26

Quote:

Oddly enough a column of molecular sieves will do the same thing to propane. I have no clue how though.

kilowatt - 21-8-2008 at 05:15

Zeolite is a catalyst too. It is used specifically in some catalytic cracking of hydrocarbons.

vulture - 21-8-2008 at 12:30

Quote:

Without concentrated sulphuric acid you will never go higher than 68%, due to the azeotrope of water and HNO3

Not entirely true. Magnesium nitrate and similar (compatible to HNO3 ofcourse) desiccants will break the azeotrope. It works most favorably with the DCM extraction method. They have the added bonus of being more easily recycled, plus you can make magnesium nitrate easily from the carbonate and dilute nitric acid.

kilowatt - 21-8-2008 at 17:16

Quote:

They have the added bonus of being more easily recycled, plus you can make magnesium nitrate easily from the carbonate and dilute nitric acid.

That forms magnesium nitrate hexahydrate though, which is not as strong a dessicant. It cannot be dehydrated by heating alone; it will just decompose to magnesium oxide and NO2/O2/H2O vapors. How can it be converted to anhydrous magnesium nitrate besides with anhydrous HNO3?

Rosco Bodine - 21-8-2008 at 18:50

Magnesium nitrate doesn't have to be dehydrated completely to have sufficient dehydrating power to break the azeotrope of nitric acid and water. IIRC, even the trihydrate will do the trick. It can only be conveniently obtained anhydrous by heating alone AFAIK, in double salt form with NH4NO3. The patents give good information on this.

Contrabasso - 22-8-2008 at 10:00

Thanks all, 98% H2SO4 is available from a few suppliers here but nitric rarely exceeds 50%. SO H2SO4 + 50% HNO3 will distil off 98% HNO3 then a largely water portion will come over leaving conc H2SO4

Looks like a simple distillation taking the water fraction out in the middle

kilowatt - 22-8-2008 at 10:21

As long as by "simple distillation" you mean fairly easy, but with a reflux distillation as opposed to a simple distillation.

Quote:

The patents give good information on this.

I beg to differ; the patents relating to anhydrous magnesium nitrate describe methods with a pulse combustion drier. Double salts with ammonium nitrate were briefly referenced but not described.

[Edited on 22-8-2008 by kilowatt]

Rosco Bodine - 22-8-2008 at 21:00

The attached patent confirms what I just said above.

You are onto some idea of baking the nitric out of a nitrate
like baking out the sulfuric from a sulfate, in order to cycle through ammonium nitrate and recover ammonia and convert it to nitric......and the energy requirement for all the requirements of that loop are economically infeasible
as well as being a technical nightmare. So I am not even
suggesting usefulness in that regard.

But if you want something to break the azeotrope,
Magnesium nitrate will do it and it doesn't need to be anhydrous. IIRC there are other patents also which
describe this usefulness and it is actually more thermally efficient than H2SO4 on the regeneration of the dessicant.

It hasn't been the source of any experiments by me to determine if it is workable, but the anhydrous double salt of magnesium nitrate and ammonium nitrate may possibly also
break the HNO3/water 68% azeotrope. This would make an interesting experiment. See the patent US3173756 anhydrous double salt of claim 11 , NH4Mg(NO3)3 mp. 167C
I have speculated that the double salt forms under only basic conditions, and separates into a simple mixture of discrete salts under acidic conditions. If my hypothesis is correct, then this would provide access to the dehydrating
potential of the anhydrous magnesium nitrate supplied as
the predominant weight percentage component of the
double salt, which is the only practical means of supplying an anhydrous magnesium nitrate component, shielded and contained within a double salt form.

http://www.sciencemadness.org/talk/viewthread.php?goto=lastp...
But by all means if you have a better idea, let's hear it

[Edited on 23-8-2008 by Rosco Bodine]

Attachment: US2463453 Magnesium nitrate for HNO3 concentration.pdf (265kB)
This file has been downloaded 785 times

kilowatt - 23-8-2008 at 11:23

Correction: the patents I was able to find on my own.

Would hygroscopic sulfates like copper sulfate also work for breaking the azeotrope?

[Edited on 23-8-2008 by kilowatt]

Rosco Bodine - 23-8-2008 at 12:06

To greater or lesser extent there is usefulness for many of the dehydratable salts, and for mixtures, eutectics and
double salts, since there are combinations which dehydrate completely and at much lower temperature than
would the single salt. In some of these schemes, the dehydration works by one salt substituting for the water of hydration which would normally be bound in another.
Magnesium nitrate and calcium nitrate, zinc nitrate, ammonium nitrate are candidates. The more usual sodium and potassium nitrates can also be useful as components
in eutectic mixtures, and the lowered melting point facilitates the thermal dehydration occuring at a much
lowered temperature. Some presence of a basic salt
in small amount is probably also involved in lowering the melting point and facilitating the dehydration, accounting for the observation that the desired effect only seems to
occur in basic mixtures, the mixtures being disrupted by
acidic conditions, which can lead to the distinctly separate
hydrates crystallizing out separately as a mixture, no longer a eutectic or complex double salt, but the common hydrates appearing as a mixture of the usual hydrate crystals.

merlic79 - 25-8-2008 at 17:16

Today I produced 150ml of fuming nitric acid with a density of 1.48g/l. The distillation setup was very simple; it consisted of a portable electric stove top element, a 12L high quality stainless steel pot, and a gallon corningware glass bowl filled with ice water. In order to to collect the distillate, I suspended an glass cup wrapped in rockwool insulation and Al foil. 500g (excess) of KNO3 was added to the pot, then 250 ml of Drain Opener was added (I can get this stuff as I have a plumbing ticket). Heating commenced and the icewater bowl was put onto the pot. Fumes were seen condensing onto the cold glass, and these droplets fell into the insulated
cup, after 1h heat was removed and the nitric acid was measured. I plan to water down this acid, as I do not need it so strong (down to 70% for me). I intend to now make some Mono Nitrotoluene, and then reduce to P-toluidine using Sodium dithionite (Rust Out)........has anyone tried this reduction before?

h2o2guru - 26-8-2008 at 08:50

A high temperature arc discharge will produce NO, a corona discharge will produce ozone.

There is a wonderful science experiment you never see anymore. that's running a Jacobs ladder in a clear plastic tube, after a few minutes you can't see thrue the tube anymore, from the NO gas.

This was actually a method of production of nitric acid and nitrates at the turn of the 20th century. What I find interesting is that its a more efficient process than whats used now, but uses a more expensive form of energy (electricity)

I'm presently putting together a system that will give me All the nitric acid and nitrates I might ever want !

when I'm closer to finishing I'll post some photo's

Contrabasso - 26-8-2008 at 09:48

A process for Nitric acid would be most helpful and fit alongside the Lead Chamber process described for the manufacture of Sulphuric acid.

franklyn - 26-8-2008 at 17:39

I've thought about this some and I believe 68% nitric acid can be dried prior
to distillation using Chile saltpeter and Epsom salt in the following way.
One needs first to dry the MgSO4•7H2O by cooking in an oven at high heat
for a couple of hours. The proceedure is to mix into the acid both salts which
metathetically reform in solution as hydrated salts.
* Note that more water is indicated below at the right than on the left
so that the hydrated forms shown will not be realized in practice.
Na2SO4•7H2O and Mg(NO3)2•2H2O are the lesser hydrated variants.

. . . . . . . . . . . 17 HNO3 + 28 H2O | + | 4 NaNO3 + 2 MgSO4 => 2 Na2SO4•10H2O + 2 Mg(NO3)2•6H2O
. . . . . . . . . . . . . . . . | . . . . . . . . .| . . . . . . . . . . .| . . . . . . . . . . | . . . . . . . . . . . . | . . . . . . . . . . . . . . . . . . |
molar mass . . 1071 . . . . . 504 . . . . . . . 340 . . . . . . . .240 . . . . . . . . . .644 . . . . . . . . . . . . . . . 512

1575 grams of 68 % acid are represented above which may be proportioned for
any lesser amount. So for 500 grams of acid at 30 ºC the proceedure would be
to divide it into two beakers , one holding 300 grams of acid and the other 200.
Add and mix in 108 grams of NaNO3 to the beaker with 200 grams acid, then add
and mix in 76 grams of MgSO4 into the other beaker with 300 grams acid.
One then pours and stirs both amounts togther , promptly placing into the freezer
for about half an hour to chill. Solubility of Sodium sulfate drops from a maximum
of 49.7 grams per 100 g water at 32.4 °C to almost nothing at 0 °C .
In the presence of other more soluble salts the solubility of Sodium sulfate is
markedly diminished. Hence when the mixture is retrieved from the freezer,
a precipitate of just about 204 grams of hydated Sodium sulfate should be plainly
seen. After filtering the now much more concentrated acid Mg(NO3)2 mixture is
ready for distillation according the patent given above US2463453.

Solubilities at 30 ºC
- Source - Solubilities of Inorganic and Organic Substances
MgSO4 - 29 % of solution or approximately 409 grams per liter H2O - page 184
NaNO3 - ~49 % of solution or about 960 grams per liter H2O - page 307
*note: what is indicated for NaNO3 in the reference is more than what is
calculated by using the 8.o6 mols per liter given.

Related posts by garage chemist
http://www.sciencemadness.org/talk/viewthread.php?tid=9450#p...
http://www.sciencemadness.org/talk/viewthread.php?tid=1851#p...
http://www.sciencemadness.org/talk/viewthread.php?tid=2823#p...
and one from Bromic Acid
http://www.sciencemadness.org/talk/viewthread.php?tid=946&am...

.

chief - 21-9-2008 at 09:12

I never did it -- but: Ba(NO3)2 + 2 H2SO4 ==> BaSO4 + 2 HNO3
BaSO4 should ppt., ready ! Has anyone tried this ??

The Ba(NO3)2 might stem from fireworks (those brillant-candles, on the iron-wire):
==> ppt out BaCO3
==> react with some dilute HNO3
==> dry the Ba(NO3)2
==> do the conc. HNO3-reaction, yielding BaSO4

--> recycle the BaSO4, using soda (there was a thread on this, by me)

Raw BaSO4 may be found as filler-material in heavy-weight high-polish-paper !!
Those -heavy- books contain BaSO4 as filler !

Besides: The alchemists, as I believe, distilled saltpeter with Iron- or Copper-vitriole; the forming H2SO4 reacted with the saltpeter, etc . (Copper-vitriole + "Alaun" + Saltpeter: http://www.retrobibliothek.de/retrobib/seite.html?id=114115)

[Edited on 21-9-2008 by chief]

chloric1 - 21-9-2008 at 10:38

Nice franklyn but we are not dissolving your salts in pure water we are dissolving them in concentrated HNO3. I am not so sure you will get mcuh of the salts to dissolve. Sodium nitrate might be only slightly soluble in concentrated HNO3 because of common ion effect. Secondly, dried epson salts might not dissolve at all.

Xenoid - 21-9-2008 at 11:53

Quote:

Besides: The alchemists, as I believe, distilled saltpeter with Iron- or Copper-vitriole; the forming H2SO4 reacted with the saltpeter, etc . (Copper-vitriole + "Alaun" + Saltpeter

Yes, I have often wondered why the old method attributed to "Geber" has not been mentioned in this thread; distillation of copperas (ferrous sulphate) with saltpetre and alum. All the ingredients are available from a garden centre.

Does anyone know the exact function of alum in this procedure?

Formatik - 21-9-2008 at 12:39

Quote:

Originally posted by chief Besides: The alchemists, as I believe, distilled saltpeter with Iron- or Copper-vitriole; the forming H2SO4 reacted with the saltpeter, etc . (Copper-vitriole + "Alaun" + Saltpeter: http://www.retrobibliothek.de/retrobib/seite.html?id=114115)

Yeah it works. Heating on low flame CuSO4.5 H2O with KNO3, the mixture turns green and gives off red-brown NO2 gases. I never tried to condense it though. Heating Cu(NO3)2 itself at low heat will already form NO2 gases. Maybe it could form and decompose in situ by heating KNO3 with some neutral or acidic inorganic copper salts.

[Edited on 21-9-2008 by Formatik]

chief - 22-9-2008 at 04:38

Instead of Ba(NO3)2 + H2SO4 ==> 2 HNO3 + BaSO4
possibly also : Pb(NO3)2 + H2SO4 ...
might work, and Pb is more available.

PbSO4 is as insoluble as BaSO4, and (I _believe_) also insoluble in conc. HNO3

But no guarantee, think yourselves ... !

crazyboy - 13-11-2008 at 17:41

Do you think this vacuum can pull sufficient pressure to distill nitric acid with no/minimum decomp? http://www.surpluscenter.com/item.asp?UID=2008111217250165&a...

watson.fawkes - 13-11-2008 at 18:17

Quote:

Originally posted by crazyboy
Do you think this vacuum can pull sufficient pressure to distill nitric acid with no/minimum decomp? http://www.surpluscenter.com/item.asp?UID=2008111217250165&a...

Oh no. Not even close. That pump is specified as 25" Hg max. That's 635 mm Hg. Subtracting that from 1 atm = 760 mm Hg, you get 125 mm Hg. I'm not sure that's even classified as vacuum, just low pressure.

S.C. Wack - 13-11-2008 at 19:37

Well, that's enough vacuum to lower the bp of water to 43C...I sometimes use an oilless Gast that pulls that (not for this), because you don't always need a turbovac. I think a relevant question here is what, exactly, is a reasonable maximum pot temperature for doing this.

Though a Nalgene aspirator should really be the way to go. I would not trust any train of base to clean it up completely before hitting metal, though this may be colored by my intense dislike of the fuming acid.

crazyboy - 13-11-2008 at 20:23

Thanks... any suggestions? I don't want an aspirator (I have one) because I cant run a hose and I don't have a sink were I work. I tried a 1/2 HP pump but it got hot and water went back into the vacuum outlet.

By the way will this http://www.grainger.com/Grainger/items/5Z667?cm_mmc=Google%2... or this http://www.harborfreight.com/cpi/ctaf/displayitem.taf?Itemnu... work?
I'm a bit confused with its vacuum capabilities maybe someone has some idea?

Also and more importantly I made the poor choice of buying a threaded thermometer adapter and my o-ring and cap got melted by nitric acid even though I wrapped them in Teflon tape. I know I need a side arm adapter with three ground glass joints but I cant seem to find a 24/40 thermometer adapter that will hold a vacuum and not get corroded by hot nitric acid.

Please help.

[Edited on 13-11-2008 by crazyboy]

ppoowweerr - 14-1-2009 at 04:43

I was playing around the net and found a "Nitric acid substitute". I couldn't find an MSDS or anything that was informative. It said that the ONLY time this salt was a suitable replacement for nitric acid is when added to HCl ie. to make aqua regia. Does anyone know what this "salt" is? The only thing that comes to mind is ammonium nitrate, but since I have never combined the two materials I am not sure if it is viable.

I dont really need the pure nitric acid since I don't want to nitrate any energetics, aqua regia is my intended product so any information is appreciated.

not_important - 14-1-2009 at 20:55

Likely sodium or potassium nitrate, or possibly a persalt that oxidises the HCl to free chlorine.

Contrabasso - 15-1-2009 at 11:30

Vacuum!
I have seen a scheme whereby an aspirator is fed by a waterpump recirculating a reasonable volume of water. Now I really fancy this with a plastic aspirator and a plastic pump all the corrosive waste should be in the water for sensible disposal, AND the pump should live reasonably long (unlike an oil filled rotor pump).

hissingnoise - 16-1-2009 at 06:28

Adding small amounts of NaHCO3 to the pump-water might protect the pump from acids carried over by the vacuum.

hissingnoise - 16-1-2009 at 07:30

Then again, IMO, the increase in HNO3 concentration afforded by vacuum-distillation really isn't worth the extra hassle.
Atmospheric distillation from anhydrous (predried) KNO3 and 98% H2SO4 gives RFNA which can be decolourised by bubbling ozonised air or oxygen through.
~98% HNO3 can be prepared this way without resorting to (awkward) vacuum set-ups.

3287 - 16-1-2009 at 09:19

It's been mentioned that platinum and copper can be used to oxidize ammonia.

Transition metals in general are often used for purposes of oxidation. Could another transition metal catalyst be used instead, making the process more suitable for the labless mad scientist?

hissingnoise - 16-1-2009 at 10:18

The oxidation of ammonia on copper does not produce NO---NH3 is decomposed by the reaction.

Bohrium - 16-1-2009 at 12:21

Nitrates do not react as 2XNO3 + H2SO4 --> X2SO4 + 2HNO3.

Because H2SO4 is diprotic, it will react like this: XNO3 + H2SO4 --> XHSO4 + HNO3.

hissingnoise - 16-1-2009 at 13:28

Quote:

Originally posted by Bohrium

Because H2SO4 is diprotic, it will react like this: XNO3 + H2SO4 --> XHSO4 + HNO3.

Actually both protons can be used as in this scheme; 2KNO3 +H2SO4---> K2SO4 + 2HNO3.
A smaller quantity of H2SO4 is used but the bisulphate is the preferred byproduct in practice. . .

vulture - 16-1-2009 at 13:45

Quote:

Actually both protons can be used as in this scheme

No they can't, HNO3 is a stronger acid than HSO4-.

Jor - 16-1-2009 at 13:49

Are you sure vulture?

At room temperature yes, then only the first proton can protonate KNO3 or similar, as the HNO3 is not yet a gas at those temperatures.

But at higher temperatures, above the boiling point, the equilibrium is pushed more and more to the right, because HNO3 escapes from the system as a gas.

It is also possible to produce HCl from NaHSO4....

hissingnoise - 16-1-2009 at 14:05

With respect, vulture, it's to do with the volatility of the acids, the pot-temperature and the stoichiometry of the reactants, not acid-strength. . .

Bohrium - 16-1-2009 at 20:39

Quote:

Originally posted by hissingnoise
With respect, vulture, it's to do with the volatility of the acids, the pot-temperature and the stoichiometry of the reactants, not acid-strength. . .

HSO4- will not disassociate in any considerable yield, so it will be unable to bond to free NO3- and boil off. So, you will be left with a bunch of HSO4- and X+ in solution, which will eventually crystallize.

[Edited on 16-1-2009 by Bohrium]

hissingnoise - 17-1-2009 at 04:51

Bohrium, it's quite simple, the proton in a hydrogen-sulphate is acidic and being acidic it is replaceable/reactive---this is elementary text-book stuff!

woelen - 17-1-2009 at 05:18

Yes, you can make HNO3 from KNO3 and KHSO4 (or the sodium salts). This is not really easy though, less so than making HCl from NaCl and KHSO4 (or NaHSO4). This is because HNO3 tends to decompose at the temperatures involved and so there will be a lot of NOx, O2 and H2O as well.

jimwig - 19-1-2009 at 09:54

me thinks a vacuum distillation will result in a high concentraqtion of hno3.

hissingnoise - 19-1-2009 at 10:20

You could, theoretically, get ~100% HNO3 by vacuum dist..
And from atmospheric dist. ~98%, again theoretically. . .
Is two percentage points worth all that extra trouble?

I wouldn't take everything Fester says in KIE as gospel!

[Edited on 19-1-2009 by hissingnoise]

vulture - 19-1-2009 at 13:19

Quote:

With respect, vulture, it's to do with the volatility of the acids, the pot-temperature and the stoichiometry of the reactants, not acid-strength. . .

With respect, unless otherwise indicated, one assumes conditions are room temperature and 1atm. In that case HSO4- will not protonate NO3- to any significant extent.

Sand also oxidizes carbon, but only at extremely high temperatures. What I'm saying is that you should give the conditions if they deviate from SOP.

[Edited on 19-1-2009 by vulture]

hissingnoise - 19-1-2009 at 13:40

Point accepted, vulture. . .my contention was, though, based on the conditions pertaining to the reaction(s) specified in the thread.
In which case, we're both correct!

Leander - 27-2-2009 at 14:21

Yesterday I made some nitric acid with my improvised distillation setup. It involves a glass beaker covered with a thick acid resistant plastic sheet folded into an upside down cone, and filled with cold water. In the center of the beaker a small glass yar is placed to collect any acid condensing on the cold sheet. The device is loaded with 150 grams of H2SO4 (96%) and 100 grams of KNO3.

I distilled for 8 hours. First I slowly started heating until about one drop per second came over. Since the colour of the acid was only slightly yellow I decided to manage that temperature, probably 100-150oC.

The yield was 38,8 ml. The density of the acid was measured using using a pipette and a digital scale with an accuracy of 00.1 gram. The density was about 1.54 gr/cc, while the maximum density of 100% HNO3 is 1.513.

I wondered if it's possible that some H2SO4 came over? Has anyone ever had this problem before? I'll check this in a couple of days using some BaCO3.

And how is it even possible that I yielded only slightly yellow acid, while the stuff should have been spewing NOx at that temperatures?

[Edited on 27-2-2009 by Leander]

hissingnoise - 27-2-2009 at 15:43

Better, IMO, to use a proper distillation set-up for HNO3.
If the plastic cone wasn't teflon, or similar, some coating on its surface may have dissolved in, and reacted with your acid---it's also just possible that H2SO4 splashed into the "distillate".
It would explain the highish density. . .
NO2 in HNO3 also increases the reading and it lightens in colour somwhat as the temperature is reduced.
In any case, acid produced in this way probably shouldn't (for safety reasons) be used to prepare compounds like, say, cyclonite.
You need just an RBF, condenser and receiver---simple and inexpensive.
Eight hours for 39cc. of pretty dodgy HNO3, Jeezzz!

Leander - 28-2-2009 at 08:37

The plastic cone became slightly white after the exposure to nitric acid. No serious decrease in thickness was observed.

Further, I wasn't planning to use this in any RDX synthesis. But why should it be dangerous to do so anyway?

If gives me great satisfaction to design and use improvised apparatus and synthesis. IMO is't not really a challenge, or as much fun to obtain al sorts of chemicals if I was equiped with a full proffesional lab. Just a matter of interest, has nothing to do with k3wlyness of lack of equipment or chemicals.

Aqua_Fortis_100% - 28-2-2009 at 09:22

@hissingnoise , PEAD sheet (Ive used supermarket package) will stand relatively well for some uses before that you need to replace it; and just to be on safe side, you can use it one time then replace it.. Ive used this method for producing my nitric acid, and is the most inexpensive and OTC that I came across. All you need for glassware/setup is a ultracheap otc jar , plastic and a moderate heat source..

Quote:

Originally posted by Leander
If gives me great satisfaction to design and use improvised apparatus and synthesis. IMO is't not really a challenge, or as much fun to obtain al sorts of chemicals if I was equiped with a full proffesional lab. Just a matter of interest, has nothing to do with k3wlyness of lack of equipment or chemicals.

That is also what I always have wondered and what inspire me and is just fun to obtain hard to acquire chemicals through DIY schemes.... Just think that in a not too much distant future they watch chemical glassware selling.. Then the only way to go is improvise all.. In this case, even improvising their H2SO4 and nitrate, as some persons want..

hissingnoise - 28-2-2009 at 09:36

Strong nitric acid with organic contaminants can be dangerously explosive.
An improvised aluminium (not foil) condenser would be safer.
Near ideal would be an aluminium container with a large convex dent on its underside to facilitate dripping.
The big drawback would be the high temperature of the receiver---very inefficient!
Attempted nitrolysis of hexamine using H2SO4 contaminated HNO3 could cause a runaway because hexamine is attacked by H2SO4. . .
IMO, watching HNO3 come over at a rate of several drips ps is fun, too.
With proper distillation, you'll know your acid is pure enough for most porpoises. . .

chief - 28-2-2009 at 10:17

I just wonder, how everyone here takes a lot of trouble distilling the HNO3, still after I posted maybe already 3 times the possibility below, here a excerpt-copy of my post one page before:
=================================================================

I never did it -- but: Ba(NO3)2 + 2 H2SO4 ==> BaSO4 + 2 HNO3
BaSO4 should ppt., ready ! Has anyone tried this ??

The Ba(NO3)2 might stem from fireworks (those brillant-candles, on the iron-wire):
==> ppt out BaCO3
==> react with some dilute HNO3
==> dry the Ba(NO3)2
==> do the conc. HNO3-reaction, yielding BaSO4

--> recycle the BaSO4, using soda (there was a thread on this, by me)
============================================================

If this works (and I remember to have read about it in some chemistry-book too), then its as simple getting HNO3 + Ba(NO3)2, and letting set. The 100 % HNO3 would be the liquid afterwards, the H2SO4 gone into the unsoluble and heavy BaSO4 (setting fast) ...

Besides: The Ba(NO3)2 does not decompose below 595 Celsius, thats red-glowing. It can therefore be dried efficiently. With the nitrate so dry, and a dry H2SO4, there seems to be no reason for any much water in the HNO3 later on ...

hissingnoise - 28-2-2009 at 10:18

Quote:

Originally posted by Aqua_Fortis_100%
Just think that in a not too much distant future they watch chemical glassware selling.

Those clouds on the horizon are nearly directly overheard (Bushwhackered?).
Get yer flasks'n'condensers while you still can. . .
Sorry, where was I---oh yeah, I used the field-expedient method, too.
Just once. . .

Aqua_Fortis_100% - 28-2-2009 at 10:35

Yeah, but with nitric made just one time with PEAD I dont think that organic contaminants will be a great issue, and since most(?) of them are probably oxidized to CO2/H2O, is safe.. Anyway for me at least there is even added safety since I never used that nitric for pyro purposes, just to make inorganic nitrates that are good to keep around; I would prefer plain H2SO4/KNO3 nitrating system , although it cant be used to nitrate things like hexamine and can be difficult to others like glycerin.. The major hazard that always worried me are NOx fumes, both in nitric synth and nitration experiments.. Very nasty thing.. You can detect by smell it in very minute amounts, but like many substances on prolonged exposure you will be no longer able to detect its acrid smell... They make your lungs go into soup , even weeks after exposure (very acute exposure = immediate pulmonary edema/death) ; and since your lungs have few nerves you cant feel anything (mild pain in the chest, at most) until its too late.. Actually NOx fear is the only reason that I have to refuse HNO3 synth with improvised apparatus everytime I can avoid it..

I agree that is very difficult to prepare white nice crystal clear HNO3 free from dissolved NOx only with distillation ; i.e. without urea nitrate/dried warm air trick .. It was mentioned here, IIRC, that even minute exposure to sunlight will make HNO3 releasing some NOx and therefore getting yellow tint..

hissingnoise - 28-2-2009 at 10:47

Quote:

Originally posted by chief
I never did it -- but: Ba(NO3)2 + 2 H2SO4 ==> BaSO4 + 2 HNO3
BaSO4 should ppt., ready ! Has anyone tried this ??

Chief, if you add dry barium, lead or calcium nitrate to H2SO4 the insoluble sulphate will form on the outside of the granules preventing further reaction.
Some small amount of HNO3 may form but most will be locked within the grains.
Normally, a saturated stoichiometric solution of the nitrate salt is used, giving dilute HNO3 and the sulphate precipitate.

Aqua_Fortis_100% - 28-2-2009 at 10:51

@ chief
Ba(NO3)2 is somewhat expensive to use in HNO3 manufacturing; I would keep it for green stars... Just use Ca(NO3)2 instead, although very difficult to dry since it is deliquescent IIRC..
How easily dry Ba(NO3)2 and conc H2SO4 would react to form HNO3? Since BaSO4 is very insoluble, I dont think it be easy to get all Ba(NO3)2 reacting...
But its interesting the possibility that NaOH can digest BaSO4, so can it??? Im not sure if this really works; just a feeling though since BaSO4 is very stable, insoluble and thus cant be easily decomposed.. I just knew that only way to recover Ba from BaSO4 is through roasting with reducing agent (charcoal, etc) that can be boring, messy and time consuming or then making a homologue of "incendiary plaster" with BaSO4/aluminium just for eyes amusement.. Then getting the BaS and dissolving in HCl, avoiding toxic H2S and then converting chloride to nitrate through NH4NO3 (since Ba nitrate has low solubility on low temp.. Ive made this before) or then making BaCO3 / Ba(OH)2 from aqueous chloride... Overall, messy process.. But its still an alternative...

EDIT:

Quote:

Originally posted by hissingnoise
Chief, if you add dry barium, lead or calcium nitrate to H2SO4 the insoluble sulphate will form on the outside of the granules preventing further reaction.
Some small amount of HNO3 may form but most will be locked within the grains.
Normally, a saturated stoichiometric solution of the nitrate salt is used, giving dilute HNO3 and the sulphate precipitate.

Oh, sorry.. Ive posted almost the same thing, you beat me first ahaha

[Edited on 28-2-2009 by Aqua_Fortis_100%]

[Edited on 28-2-2009 by Aqua_Fortis_100%]

chief - 28-2-2009 at 16:01

The surface-covering would be probably no problem for the following reasons:
==> SO4 is divalent, and tetrahedron
==> NO3 is monovalent and planar
===> give crystal-chemically different structures, probably uncapable of sticking one on the other

Also: H2SO4 (hot & conc.) _dissolves_ BaSO4 to some amount (as literature says), so the mixture of H2SO4/Ba(NO3)2 would only have to be heated, at most. ..

Also the reaction of BaSO4 with NaOH works in the melt, as equilibrium, but it's easier to boil it for 1 h in conc. soda-solution, thread upon this here in the forum. I didt it, and it works, is good for making Ba-compounds out of BaSO4, which could be recycled this way.

hissingnoise - 1-3-2009 at 06:02

Chief, you could make a hot saturated solution of Ba(NO3)2 and add to it the calculated amount of 98% H2SO4 with stirring---this will give you dilute HNO3 + solid BaSO4.
Decanting the solution and concentrating by distillation from its bulk of H2SO4 will give you strong HNO3.
Adding the nitrate to H2SO4 will give a solution of strong HNO3 in H2SO4 with the H2SO4 in excess.
That's how it works. . .

chief - 1-3-2009 at 13:39

Distilling only concentrates up to the azetrope.

But water-free H2SO4 + water-free BaSO4 will give HNO3 + BaSO4 ; since the BaSO4 is so insoluble, it should ppt out, and ready would be the 100 % HNO3, no water in it (except for small amounts) ; besides it's from a chembook, but forgot where to find the reference ...

Even if the HNO3 would be mixed with excess H2SO4, then this could be distilled much better, since no water in it ... !

I can't believe noone gets the idea ... , and I didn't post it for the first time . Do I miss anything ? I might as well try myself, but I'm not common to handling such conc. HNO3, also no uses for me of this stuff ... .

hissingnoise - 1-3-2009 at 14:32

Quote:

Originally posted by chief
I'm not common to handling such conc. HNO3, also no uses for me of this stuff ... .

sigh! At this point chief, there's little more I can tell you. . .

chief - 1-3-2009 at 15:47

Don't make yourself more important than you are: As you see I avoid it even though I could do it, for the chems I own; eg. the H2SO4 I own since years (and some nitrates since the 80s, most of them still unused ...), and I just could resist ... ; I'm just not a free-of-will front-pig, like most usual chemists ...

Also it's not too dangerous, if done with the proper precautions: Little amounts, face- and earprotection, thick unburnable clothing of the right sort of fibers, that can be removed without pulling it over the head, etc. . I find it lame to play such a hide-and dagger-game: "I can't tell you more at this point ... " ; anyhow someone probably is gonna try it, and will confirm, some day.

Main thing for me not to do it is: It's not really a thing I'm interested in, besides the theoretical confirmation that it works. So I post it here, where everyone anyhow is messing around with such things ; the "Im not used to handle conc. HNO3" was just only sort of a warning towards others (it isn't even really true for me), to check the dangers first, since this is a 1-minute experiment, quickly done without lot's of preparation.

Also it's probably (!) much less dangerous than having the usual conc. H2SO4+nitrate destillation at 150 [Ceslsius], where everyone who ever tried the effect of H2SO4 alone at this temp would just shy away from it ...

Also: I might do it within a minute, but I definately wouldn't go beyond that point the check if the HNO3 is "really good" or something ..., so it would be of little value if I did it: I wouldn't check the results ...

[Edited on 1-3-2009 by chief]

The Ostwald Process

Formatik - 15-3-2009 at 13:03

Despite all of the discussion on synthesis here, there has been no lengthy focus or experimentation on one of the most important processes for the preparation of nitric acid, the Ostwald Process.

I did an experiment some time ago on this where air is mixed with NH3, dried, goes through the catalyst, then into an air-mixer bottle, through water in a 1L flask, and then into a basic solution (neutralization attempt). The comments are below. A video of the set-up is here. It is a primitive set-up since these were a series of probe attempts.

1. Ammonia. The amount in the small 125mL flask is 15mL of 10% NH4OH. The air stream from the pump goes over the surface of the liquid. The explosible limits of ammonia in air are between 15 to 30%, a relatively narrow range. Assuming all of the NH3 were vaporized at once its volume would be something like 2.01mL / 110mL = 1.83%. But since the evaporation is slow, and an air stream is going through it I was working far below those limits. The receiver water in the 1L flask is about 60mL. With those amounts, maximally 3.55g HNO3 can form from: 3 NO2 + H2O = 2 HNO3 + NO. With more being possible by further combination of NO. Unwanted side reactions can occur, and these form N2 and H2O. Where sometimes there is also unreduced NH3. I didn't complete the evaporation, but reckon it takes something like 3 hrs at that amount and rate.

2. Catalyst. For the catalyst a (brass) copper net was burned in free air with a torch, then it was all pressed together. There are better catalysts to use for this, e.g. as mentioned in Gmelin, Sys. No. 4[N], 665-683. The yield of NO increases with temperature and is dependent on the catalyst used. According to Gmelin, CuO is a horrible catalyst. Tough luck. Cr2O3 will give a better conversion at a lower temp. (even active at as low as 300º, roasted pyrite also). Pure Fe2O3 will convert 9 to 9.5% pure NH3-air mixtures in good yield, optimal yield is 90% at 670 to 700º. In the industry, Pt and/or Pt-Rh is used. One could likely take a carrier like pumice, soak in a satd. solution of FeCl3, for some time and then drip in NH4OH to form Fe(OH)3, and then glow this thoroughly.

3. Drying agent and material. In the bottle above, CaCl2 pellets were used. Other drying agents will work, CaO, Na2SO4, MgSO4, etc. In the industry liquid ammonia is used, so no need for drying. Since the rubber stopper was attacked (HNO3 will attack rubber, cork, etc), I decided to wrap it in aluminium foil. It will need to be secured so no gas escapes. I also noticed a significant deformation on the glass tube after hours of heating with the bunsen burner. So something more thermally stable should be used, e.g. quartz, stoneware, stainless steel, etc.

4. Additional notes. Besides the terrible catalyst choice, air oxidizes NO sluggishly to NO2 (the rate of this conversion increases with colder temperature), this gets problematic since the stream of air is pushing this gas out. So then the NO escapes and converts to NO2 outside of entire set-up, reducing yield, posing significant and dangerous health hazard, etc. In the industry, with the mixing process with water an absorption column is used under pressure, as seen here. What I did do was lead this into a foamy NH3 solution containing a lot of NaHCO3, in hopes to neutralize it. But that didn't work and the unconverted NO would escape and soon convert to the highly poisonous NO2. Though the gases can easily be lead away using tubes.

Since then I've thought of ways to make this go more smoothly but haven't tried anything. I wonder if conc. H2O2 could be used to oxidize NO efficiently and more readily: 2 NO + 3 H2O2 = 2 HNO3 + 2 H2O

As well as combining with any NO2:

2 NO2 + H2O2 = 2 HNO3

ostwald.png - 224kB

kilowatt - 17-3-2009 at 18:40

I just rigged up an NO2 producing Jacob's ladder, since I had all the parts around from previous acquisitions. The idea is to run it continuously and produce some amount of dilute acid over days or weeks of operation, then distill it afterward. It has oiled cooled copper rails (each consisting of a tube within a larger tube and coaxial flow) and a water cooled glass jacket. Air is fed through the glass vessel with an aquarium air pump and then bubbled through water to make the acid. The ladder runs on a 7200V 30mA neon sign transformer and has a central Gabriel electrode to improve starting.

I am having some leakage problems with the oil pump (a hydraulic gear pump which I have coupled to a 19rpm high torque gear motor); the shaft seal is leaking, and I will need to turn it upside down to lessen the problem. The motor also struggles to push the oil through the narrow circuit, and will stall unless the oil is warmed up prior to turning it on. A fully integrated, preferably submersible, pump would be preferable but I don't have one.

The NO2 production rate is quite significant; within a minute of running the red color can clearly be seen inside the envelope, and the distinctive smell can be noticed quite strongly at the outlet tube.

hissingnoise - 18-3-2009 at 07:46

Very nice set-up, kilowatt: I once had something similar and found that cooling wasn't really necessary.
I used aluminium electrodes as I assumed copper would react and the cheap soda-glass jar I used got hot but not too hot to touch.
I used two NSTs in parallel and the vapour in the jar was quite dark in colour.
I figured I needed some kind of packed absorption tower to prevent losses but I never got around to making one.
I did get some very dilute acid but a tower would have made all the difference.
I will try it again now that I've got a good compressor; the original was a fridge-compressor and worked poorly.
A fully maintenance-free Jacob's Ladder would be a continuous source of acid that doesn't use-up H2SO4 for concentration, unlike nitrates. . .

kilowatt - 18-3-2009 at 09:04

Perhaps some air circulation will be sufficient to cool the electrodes? They will definitely need some cooling for continuous duty, in my experience they get *very* hot. The cooler are they less they will react too, but I have found that copper is not very reactive anyway at least over the time periods and concentrations I have run this so far.

[Edited on 18-3-2009 by kilowatt]

watson.fawkes - 18-3-2009 at 11:00

I might suggest a pair of tungsten rods for TIG welding for electrodes in such an arc reactor. The ceriated and thoriated versions have lower work functions, so it's easier to sustain the arc. As a refractory metal, tungsten rods shouldn't need cooling, as long as the contact point of the arc moves as it should.

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