Sciencemadness Discussion Board - help make white fuming nitric acid.

help make white fuming nitric acid.

Pages: 1 2

avi66 - 26-4-2010 at 03:36

hello, i build a vacuum distillation system as shown on the picture below, i run this setup with oil temperature 140-120c, 210 gram kno3 and 106 gram h2so4, i run it for 1.5 hour, the thermometer didn't change its was 28c, my lie-big condenser was 30cm long and already get water from ice-water bate, i got alot of nitrogen oxides, and my yield was poor(about few drops), the nitrogen oxides bubble alot, at my ca(oh)2 filter, and i see alot of redish vapor in 2L boiling flask and circulating white fumes inside it, i think because of my ca(oh)2 filter my system build little bit preasure 1-1.2 bar, and all my hno3 fumes heavy so soo they circulate in 2L flask until they became oxides+water, is my theory true ?
i also try remove my bubbler , to allow the nitrogen oxides go rapidly to atmosphere and let my hno3 fumes climb to the condenser, i think its help little, is it possible that the big size and spherical shape of my 2L flask cause the circulation+decomposition of my pure hno3?
i see in my setup that the bubbler don't work as good as i thought, i was able to see above my calcium hydroxide solution red fumes, which are dangerous for my water aspirator, anyone have an idea how could i make trap for the nitrogen oxides, i cant get solid carbon dioxide.
and anyone know how i can help my hno3 fumes climb fast to my condenser?
Thanks for all who pay attention !

hissingnoise - 26-4-2010 at 04:25

Use equal weights of nitrate and H2SO4 and forget about vacuum distillation.
Your condenser and receiver aren't cold cold enough to condense the HNO3 vapours.
Receiver temperatures close that of dry-ice are needed to condense HNO3 vapours at aspirator pressures.
Without vacuum your HNO3 will be coloured by NO2 but this can be removed to a large extent by blowing dry air through it.
If you need water-white HNO3 just add a small pinch of urea as a last resort.
The bubbling in your filter, BTW, was caused by HNO3 vapours reaching the lime and being neutralised.

quicksilver - 26-4-2010 at 05:10

Amen.
I have distilled quite a bit of HNO3 over the course of years and found certain truths. You really don't need the vacuum.
The "sweet spot" of temp is 80 C - and a distinctly cold reception vessel. There is where you'll get some serious flow.

avi66 - 26-4-2010 at 06:12

as i write above, i didn't use my vacuum pump in the experiment.
red fuming nitric acid is suitable for nitroglycerin synthesis?
so, if i use same weight of kno3 and h2so4 il spend valuable sulfuric acid, i read that i can use as twice kno3 by more heat:
http://www.lateralscience.co.uk/1888chem/experiments.html (look at the middle of the page)
so you say that without dry ice nitric acid vacuum distillation isn't practical look at this video, he didn't talk about dry ice:
http://www.youtube.com/watch?v=CtdX5YmOdcs

and this one make the acid without long condenser that i have:
http://www.youtube.com/NurdRage#p/u/48/2yE7v4wkuZU

is he use salt-water solution ? how he got batter results? salt-water solution can really help the condensation?

i see another man which use water-ice bath and successfully make hno3 with vacuum distillation:
http://www.craigsarea.com/hno3.html
how you can explain that he condense the product with only salt ice condenser under vacuum?

The WiZard is In - 26-4-2010 at 09:54

See Nitric Acid in Brauer - Handbook of Preparative Inorganic Chemistry
p. 491. La book is in the library.

franklyn - 26-4-2010 at 12:35

hissingnoise & quicksilver succintly stated the process shown in Brauer
Not having done this myself I only relate here what is written in older literature.
HNO3 is prepared by distillation of Sodium Nitrate with concentrated Sulfuric acid
in a cast iron retort forming Sodium Bisulfate and condensing the Nitric acid vapor
in a cold receiver. This is not easily done well without understanding the process
procedes in two main stages and the transitions in distillation along the way.
Distillation readily occurs just over 86 ºC , raising the temperature causes distilled
vapors to decompose thus , 4 HNO3 => 4 NO2 + 2 H2O + O2 , producing Oxygen
and Nitrogen dioxide which dissolves in the initial Nitric acid becoming yellow tinged
fuming Nitric acid.
Above 200 ºC Sodium Bisulfate further reacts with more Sodium Nitrate , becoming
Sulfate , liberating all the Nitric acid only as the temperature rises to 250 ºC.

If Sulfuric acid is used in excess , the highest concentration of Nitric acid distills
at lower temperatures becoming more diluted by water and even Sulfuric acid as
the temperature rises. Nitric acid is initially more concentrated because the water
is fixed by Sodium Bisulfate forming a hydrate which decomposes at 150 ºC.
If limited to the lower temperature , less Sulfuric acid is used than needed for
complete formation to Bisulfate to ensure the retention of water. To obtain fuming
Nitric acid , 2 mols of nitrate to 1 mol of Sulfuric acid is used for Sulfate to form as
the temperature elevates. See page 383, 385, 400, 401
A Treatise on General and Industrial Inorganic Chemistry
http://www.sciencemadness.org/talk/viewthread.php?tid=6664&a...

( Avoiding decomposition with use of HNO3 also has peculiarities that must be observed )
From footnote 1 on page 400 in the above reference _
" Whereas nitrous gases are readily expelled from nitric acid, they are not easily eleiminated
from the concentrated nitric-sulphuric acid mixtures commonly used in explosive factories,
and it is best to prevent their formation during the mixing of the striong nitric acid with the
sulphuric acid by cooling the mixing vessel outside with a spray of water and taking the pre-
caution to add the sulphuric acid to the nitric acid (not vice versa) and stir the mass well."
* My note * This is particularly troubling if nitration is done directly in the H2SO4 - nitrate salt
mixture, requiring extensive cooling , even though brown fumes inevitably result.

( The order in which ingredients are mixed can effect considerably the product yielded.)
Excerpted from page 2-14 of Military Explosives TM 9-1300-214
" During an inspection of a small Canadian TNT plant at Beloeil near Montreal in 1941,
LTC John P. Harris of Ordnance discovered that the plant was "doing things backward"
by putting toluene into the acid instead of putting acid into the toluene. Despite some
resistance by US TNT producers, the new process was tried at the partly built Keystone
Plant at Meadville, PA. The result was a trippling of TNT output. Lines designed to turn
out 16 tons a day produced more than 50 tons a day. The need for TNT substitutes
vanished, and the cost per unit was cut in half."

.

[Edited on 27-4-2010 by franklyn]

The WiZard is In - 26-4-2010 at 13:07

Quote: Originally posted by franklyn

hissingnoise & quicksilver succintly stated the process shown in Brauer
.

There am being a large difference in usefulness — between stating and referencing.

Chemical engineering is in the details. A drawing or two helps.

Prudent
Proper
Planning
Prevents
Piss-poor
Performance

The WiZard is In - 26-4-2010 at 13:20

Quote: Originally posted by franklyn

hissingnoise & quicksilver succintly stated the process shown in Brauer
Not having done this myself I only relate here what is written in older literature.
.

I recommend :—

Allin Cottrell
Nitric Acid and Nitrates
Gurney and Jackson - Edinburgh 1923
454 pages.

franklyn - 26-4-2010 at 19:00

Full title of citation here _ (" There am being a large difference in usefulness — between stating and referencing.")
http://books.google.com/books?id=Oz5DAAAAIAAJ
Not available in public domain , search for repository near you here _
http://www.worldcat.org/title/manufacture-of-nitric-acid-and...

.

Microtek - 27-4-2010 at 01:50

I have never had problems condensing nitric acid in a vacuum distillation, using simple tapwater in the condenser. You just have to include some way to regulate the pressure in the system. I add a T-section between the pump and the rest of the system. A short length of rubber hose fitted with a screw clamp is attached to the free branch of the T-section, and this is then used to regulate the pressure.
The screw clamp is then adjusted until nitric acid comes over at 50-60 C.

When the distillation is done, the screw clamp is slowly opened before the pump is turned off. This makes it easier to shut the operation down without suckbacks or other nasty surprises.

Edit: I should just clarify that I never distil nitric from sulfuric and a nitrate, but rather from sulfuric and dilute nitric. It is easier for me to buy dilute nitric acid than sodium- potassium- or ammonium nitrates.

[Edited on 27-4-2010 by Microtek]

hissingnoise - 27-4-2010 at 02:25

Variations in water pressure would make an aspirator less easy to control than a vacuum pump.
But even at 60*C I'd expect some dissolved NO2. . .
Brauer uses 36-38*C!

avi66 - 27-4-2010 at 04:08

i can control the water pressure my aspirator get with a valve, if il make 100g h2so4 + 100g kno3 in 0.1-0.5 bar pressure, using a water-ice bath for condenser i will get my fuming white nitric acid? i ask it here ... because its hard for me to run my system, i don't have time to spend.
Thanks for all who help.

[Edited on 27-4-2010 by avi66]

hissingnoise - 27-4-2010 at 08:32

You can't control the variations in water pressure which sometimes occur in municipal water-supplies. . .
But I have used RFNA in MA to prepare nitro and alone to prepare RDX without incident.
WFNA is specified for PETN and this may be due to the increased danger of runaway using RFNA.

avi66 - 27-4-2010 at 09:59

so, i read from alot of sources that vacuum distillation of nitric acid is practical with normal ice bath liquid in condenser, so if i use low temperature(50c) with 0.1 atm preasure i will get wfna(kno3 + h2so4 = khso4 + hno3) ?
here is one:
http://www.craigsarea.com/hno3.html

franklyn, its a fact that nitric acid decomposes above 86c, so how this reaction

khso4 + kno3=hno3 +k2so4) can occur at 200-250c without cause the thermal decomposition of the nitric acid?
i think il try this experiment by myself after i recrystallized my khso4 which right now dehydrate under the sun force, i will put kno3 in molten khso4, and check for nitric acid to came out.

in my lest experiment which i mentioned above, i used potassium nitrate which where slighty mixed with iron fertilizer, which make it slight pink powder(really slight), it is possible that the impurities cause the formation of hno3 to decomposed to nitrogen oxides ?

franklyn - 27-4-2010 at 10:56

Quote: Originally posted by avi66

franklyn, its a fact that nitric acid decomposes above 86c, so how this reaction

khso4 + kno3=hno3 +k2so4) can occur at 200-250c without cause the thermal decomposition of the nitric acid?

I'm not an authority I only in part quoted and referenced the source.
Concentrated acid fumes at room temperature indicating decomposition
is occuring , white if pure , red if there is nitrous acid contained. The
gas distilled is not acid per say but the equivalent of the fumes produced
when condensed and reconstituted. In actuality HNO3 is an aqueous
solution of the acid anhydride NOx ( nitrogen oxides ).

.

avi66

DetaDude - 27-4-2010 at 13:56

There are several threads on this forum dealing with this subject, by many very sharp minded chemistry people.

You may want to check out a few of these threads.

Try Our beloved nitric acid in General Chemistry (about pg. 5) this is just one on the subject there are many more excellent threads that deal with this vital acid.

franklyn - 21-5-2010 at 05:17

yada yada

NaNO3 Decomp by H2SO4.gif - 50kB

Contrabasso - 22-5-2010 at 02:34

First this is a production process for a hazardous chemical it deserves your time, attention and respect. If you can't give that don't start - you are too likely to get hurt.

To distil from dil nitric and conc sulphuric then distil under reduced pressure with a vapour temp about 60C and adjust the vacuum to hold this steadily, Be certain that the vapours condense in the middle third of the condenser -the circulating water may need cooling, Keep the product receiver cold -even iced to keep the product from boiling off!

medx - 27-5-2010 at 23:13

I used a simple distillation system to get WFNA from reaction between H2SO4 and KNO3 or NH4NO3… I made this many times.
I reached some results and I have few questions.
First results;
1- I got concentrated nitric acids. I think, at the best results, it’s concentration is about 92-5%. In some experiments product concentration was less than 90%.
2- In reaction for H2SO4/KNO3 mole ratio I used both 1mol/1mol and 1mol/2mol. If you need concentration bigger than 90% as yield of product there isn’t any utility. According theory if you use ratio of 1mol/2mol you get 2mol HNO3. But in truth after you get 1 mole nitric acid more acid decreased your acid concentration too much. For testing, after about 1 mole HNO3 occur I gathered remained acid in different container. Concentration last acid was very low. Because for second mole you must increase temperature so HNO3 decompose.
3- I reached the best concentration results when I finished reaction when I reached about 80% of theorical yield. (This means 1mol/1mol give 1mol HNO3).
4- I am not very sure but I think at first stage temperature of heater not so important. I say this for high temperature. Dropping start when temperature of reaction solution is about 85C. if temperature of heater higher it drops very fast.
These are some result which I reached. If anyone ask detail I can answer how much I know.

My questions;

Quote: Originally posted by hissingnoise

Without vacuum your HNO3 will be coloured by NO2 but this can be removed to a large extent by blowing dry air through it.
If you need water-white HNO3 just add a small pinch of urea as a last resort.

1- Does Urea change only acid color? Or does it increase acid’s concentration at same time? I added it to acid, its color changed but I think it decomposed acid. Because after this for example, acid didn’t react with hexamine.
2- What is the a pratical method for dry air?
3- What should I do to increase yield of acid which is concentration higher than 93%?
4- Are there anyone who try easier method to make WFNA?

Quote: Originally posted by hissingnoise

But I have used RFNA in MA to prepare nitro and alone to prepare RDX without incident.

I’m sorry I didn’t understand mean MA…

woelen - 28-5-2010 at 01:02

Urea does not only change color, it removes NO2 from the acid, but also makes it somewhat more dilute. Urea reacts with NO2 to form water, N2 and CO2. The gases escape from the acid, the water remains in the acid. If you have yellow acid it only contains a small amount of NO2 and then you only need a small amount of urea and the urea only makes it slightly more dilute.

If excess urea is used, then urea nitrate is formed, which is sparingly soluble. You have to add just enough urea to make the acid colorless, do not more than that amount.

Drying with air can be done with a little pump (e.g. aquarium pump) with a glass tube attached to it, which is immersed in the acid. Do not immerse plastic or rubber tubes in the acid! I personally don't like this. If you don't do this carefully you may introduce dust and other impurities in the acid and I also expect that you will have a lot of very nasty fumes of the acid going out of the bottle when air is bubbled through it. Of course you could lead the outgoing gases through another tube and immerse that in a dilute solution of NaOH to absorb the nasty fumes. The tube, however, will be eaten away quickly if this is plain plastic or rubber.

[Edited on 28-5-10 by woelen]

Jimbo Jones - 28-5-2010 at 02:18

I prefer the „purging” with dry air. The pictures below are from my homemade apparatus, but the possibilities for improvisations are endless.

http://img192.imageshack.us/i/41043.jpg/

http://img188.imageshack.us/i/41053.jpg/

http://img135.imageshack.us/i/41074.jpg/

http://img412.imageshack.us/i/41088.jpg/

For maximum results you have to add a drying chamber for the air (in this case, plastic container with calcium chloride) and hot water bath for the acid. The best part in this method (4NO2 + O2 + 2H2O => 4HNO3) is obvious, so I never tried to „clean” the acid with urea.

simply RED - 28-5-2010 at 02:35

Jimbo Jones, how hot is the water bath you use in acid purging?

Jimbo Jones - 28-5-2010 at 03:27

Around 50 – 60 °C. The starting nitric acid in the pictures (50 ml.) was absolutely clean after 40 – 45 minutes and the volume was reduced only to 46 – 47 milliliters. The same acid was used directly in the production of RDX. The yield was around 12 gr. from 20 gr. HDN.

medx - 28-5-2010 at 05:34

Jibo Jones, did you measure its concentration? How did it change?

Jimbo Jones - 28-5-2010 at 06:17

I don’t have precise scales (or desire for titration), but I’ll bet it’s over 90 % for sure. The RDX yield was given as comparison for the people who are familiar with the process.

R0b0t1 - 28-5-2010 at 12:05

I can not tell if you did or did not use your pump, but assuming you can keep any vapors or fumes from getting into and ruining the pump, I would highly suggest vacuum-distilling in the future. The resulting product would probably be water-white. (When I myself did it, the sulfuric I used had crap in it which changed the color of the resulting nitric acid... but I am assuming it would have been white)

avi66 - 28-5-2010 at 12:45

instead of make a decomposition of your fuming nitric acid using high temperature you can use batter method:
put your nitric acid under vacuum(10-30 min at 200 mmHg or 27 kPa)and the dissolved nitrogen oxides will be removed, its will limit the decomposition of your wfna ... because its will be 20-25 c instead of 60-50 c, i think 50-60 Celsius are to high for wfna ... its just make terrible decomposition.
i guess that make a hno3/h2so4 distillation to obtain wfna ... and then use vacuum to remove nitrogen oxides will give the best results for the home Chemist.

Jimbo Jones - 28-5-2010 at 13:27

“I can not tell if you did or did not use your pump, but assuming you can keep any vapors or fumes from getting into and ruining the pump, I would highly suggest vacuum-distilling in the future. The resulting product would probably be water-white. (When I myself did it, the sulfuric I used had crap in it which changed the color of the resulting nitric acid... but I am assuming it would have been white)“

I don’t really understand what you mean, but let’s try. There’s no real need of vacuum pump in the home workshop. I had very good results even with completely improvised apparatus. Lower temperatures, lower decomposition. When finish, just „clean“with dry air. By the way, even with vacuum source, the resulted nitric acid may still contain different amount of NO2.

P.S.

The „crap“in the sulfuric acid has nothing to do with your acid coloration.
------------------------------------------------------------------------------------------------------------------------------------------

„instead of make a decomposition of your fuming nitric acid using high temperature you can use batter method:
put your nitric acid under vacuum(10-30 min at 200 mmHg or 27 kPa)and the dissolved nitrogen oxides will be removed, its will limit the decomposition of your wfna ... because its will be 20-25 c instead of 60-50 c, i think 50-60 Celsius are to high for wfna ... its just make terrible decomposition.
i guess that make a hno3/h2so4 distillation to obtain wfna ... and then use vacuum to remove nitrogen oxides will give the best results for the home Chemist.“

WFNA decompose around 83 °C. Study harder.

avi66 - 28-5-2010 at 14:21

vacuum distillation is not practical because its lower the liquefied point of nitric acid ... so the hno3 need water which are more cold then ice-water mixture ....
i have vacuum pump .... how can i run get colder solution in condenser ... which allow me to use my vacuum ability?!

R0b0t1 - 30-5-2010 at 22:36

Is there a limit to the cold you can get down to (I'm not sure)? If not, mix ethanol with dry ice. You might also get a tube to go to the bottom of the vessel, making sure no fumes get sucked out by accident (depends on how leaky your setup is).

Well, yes, I can agree that there is no REAL need, but if you did happen to have one, it might be good to make use of it.

Also: I am sure the discoloration is due to what was left in the sulfic acid. I didn't actually, ehm, "distill" it. I let it stew in its own juices under a vacuum.

RFNA

The WiZard is In - 31-5-2010 at 12:48

Reminds me - a bunch and a half years ago while paging
through The Merck Index [Trivia what were the first
3 or 4 volumes of The Merck Index called?] came upon —

Nitric acid, fuming. Conc. nitric acid containing dissolved
nitrogen dioxide. May be prepared by ... or by adding a small
amount of organic reducing agent, such as formaldehyde.

What la book fails to note is that .... nothing happens for
a couple of minutes during which time you add more
formaldehyde... when suddenly you have in you Erlenmeyer
flask... a fuming nitric acid volcano, and a heck of a mess.

AndersHoveland - 20-9-2011 at 11:37

By bubbling nitrogen dioxide, along with dry air, into azeotropic (68.5% concentrated) nitric acid, the concentration can be increased up to 77%.
http://www.patentgenius.com/patent/4064221.html

quicksilver - 21-9-2011 at 12:01

This is a very interesting & useful patent. However it appears that it may only accomplish industrially due to the need for pressurized containment (8 atmospheres). I quote from the patent:

". A process for obtaining nitric acid of a concentration higher than the azeotropic concentration by means of the absorption of nitrogen oxides in diluted nitric acid,comprising the steps of:

reacting gases containing nitrogen oxides with dilute subazeotropic nitric acid to partially decompose the nitric acid by the action of NO contained in the gases, forming additional NO.sub.2 to increase the partial pressure of NO.sub.2 in thegases;

subsequently compressing the gases;

passing azeotropic nitric acid and the compressed gases containing the high partial pressure of NO.sub.2 through an absorption chamber to form super-azeotropic nitric acid;

distilling the super-azeotropic nitric acid to separate it into commercially pure nitric acid and azeotropic nitric acid;

returning the azeotropic nitric acid to the absorption chamber for use in said passing step;

injecting the gases which have been passed through said absorption chamber into a secondary absorption chamber to react with the partially decomposed dilute nitric acid from said reacting step to form sub-azeotropic nitric acid; and

returning at least a portion of said sub-azeotropic acid for use in said reacting step.

2. The method of claim 1, further comprising the step of removing water from the gases prior to said reacting step.
Description: This invention is concerned with a process for the manufacture ofnitric acid of a concentration higher than the azeotropic concentration (68% by weight) with or without the simultaneous production of nitric acid of a concentration lower than the azeotropic concentration, by means of absorption of nitrogen oxideshaving a high degree of oxidation in water or in aqueous solutions of nitric acid, for which the partial pressure of the nitrogen oxides referred to is increased partly by decomposition of nitric acid by means of nitrogen oxides having a low degree ofoxidation, and partly by compression at very high pressures of the gases that contain them.

Practically all the nitric acid produced industrially is obtained starting with ammonia by catalytic oxidation of the ammonia in accordance with the reaction:

the quantity of reagent water formed is important and when it is desired to produce acid with a high concentration it is necessary to eliminate the water almost entirely since, as is shown in the bibliography, if the water referred to is noteliminated, the maximum concentration of nitric acid possible is of the order of 77%.

The NO (nitric oxide) is considered to be a nitrogenous oxide having a low degree of oxidation, which, in the presence of oxygen and at low temperatures oxidizes to NO.sub.2 (nitrogen dioxide) of a high degree of oxidation in accordance with the reaction: - (none shown in patent)

it is precisely this nitrogen dioxide (or its dimers) that are absorbed in water to form nitric acid in accordance with the reaction: - (none shown in patent)

at a greater pressure of the nitrogen oxides that enter into contact with the water, a greater concentration of nitric acid may be obtained.

For reasons of safety, the maximum concentration of ammonia in the mixture with air for effecting reaction (I) does not usually exceed 12% if the reaction is carried out at low pressure, or 10% if its is carried out at high pressure.

The conventional processes for the production of nitric acid usually compress the gases before or after effecting reaction (I) up to the pressures of 4 or 5 atmospheres (medium pressure processes), 7 or 8 atmospheres (high pressure processes) and10 to 12 atmospheres (very high pressure). Depending on the pressures at which reactions (I) and (II) are carried out, different processes are obtained, but all of them have in common the factor that the partial pressure of the highly oxidized nitrogenoxides at the beginning of the absorption stage is not usually sufficiently high to produce large percentages of acid of a concentration higher than the azeotropic concentration (68%)."

IF this were possible in a laboratory environment it could solve many issues of acid enrichment (or "re-cycling"). Steps leading to catalytic oxidation might be dealt with (several ideas come to mind) but applications of working with NOx and common lab glass in pressure would be a tough step to take. I don't know if I would trust "Bomex" glass or similar. The heavy grade Kimax glass might hold up. If the experiment were to start with existing azotropic acid obviously this would be a non-issue.
The value is obvious but the creative bottom line is not easy as aside from scale issues the demands of a corrosive gas would entail significantly expensive materials. There MAY be a creative method of side-stepping these demands through stainless steel but superficially it seems a dedicated apparatus would be in order. Damn interesting idea though.
One question I would have is if one were to end up with RFNA and this cleaned up with Urea, would the level still hold?

AndersHoveland - 21-9-2011 at 14:06

I am fairly certain that high pressures are not required. I had also heard somewhere in the E+W forum that concentrations of nitric acid exceeding the azeotropic limit could be achieved by bubbling in nitrogen dioxide.

Pressure would not be expected to make any difference. Le Chatelier's principle only applies when an equilibrium exists. Since 98% concentrated HNO3 is relatively stable in the dark (an air tight cap can be put on it without danger of build up of pressure), and since the decomposition of highly concentrated (roughly >80%) does not seem to be reversible, there does not seem to exist any equilibrium. The pressures in the patent probably have more to do with industrial process and convenience than chemical necessity. Increased pressure means that more gas can physically dissolve, meaning faster reaction rate. If NO2 and O2 are going to react with more of the remaining water in the nitric acid solution, the gases are going to do so regardless of the pressure

I have actually tried this reaction, bubbling NO2 and air into 40% concentrated HNO3, to increase the concentration. Unfortunately, my methods were rather crude, so I cannot confirm that the concentration exceeded the azeotropic limit. But the acid definitely was much more concentrated, as indicated by comparative neutralization with bicarbonate of the prepared sample to the crontrol concentration of nitric acid.
I do not think extremely high concentrations of HNO3 can be prepared by this method, however, and the nitric acid thus prepared failed to nitrate a small sliver of paper, which had been placed into the test tube during the reaction. So 77% appears to be the limit, although I think I might have read somewhere years ago that it was something like 82% or 83.5%. ? cannot remember. I should really write these things down. For example, I once read online the precise detonation velocity of hydrazinium dinitrate, but can now no longer remember it or find it, which is very regretful. I think it was only a few hundred (about 400?) m/sec below the velocity of the mononitrate.

Are you asking if urea could be used to remove the nitrogen oxides from nitric acid? This likely would not work, since urea is a base, and would neutralize the nitric acid to urea nitrate. The NO2 would likely slowly oxidize the urea, but this would react to form more water, which would further dilute your nitric acid. I am not saying it is theoretically absolutely impossible, only that such a method would be very problematic.

[Edited on 21-9-2011 by AndersHoveland]

Reaction between NO2 and HCl to make HNO3

AndersHoveland - 21-9-2011 at 14:51

Nitric acid can also be made by reacting dry anhydrous hydrgoen chloride gas with dry nitrogen dioxide.

The HCl is dried by being passed through baked CaCl2 powder. The reactions are:

HCl + (2)NO2 --> HNO3 + NOCl

(4)HCl + (2)NO2 --> (2)H2O + NOCl

where the first is the primary reaction, and the second is the limiting reaction.

This procedure will yield fuming red nitric acid of a certain concentration, which is probably somewhere between 70-85%.

The concentration of the resultant nitric acid is probably limited by the reaction: HNO3 + 3HCl --> 2H2O + NOCl, in which nitric acid becomes a more reactive oxidizing agent than nitrogen dioxide when the solution becomes acidic enough.

Several months ago, I had a very interesting private discussion with "wiley", a member of ShadowRX forum. I would like to share that discussion here. Unfortunately, I later deleted some of wiley's posts, where showed a video of the whole setup running.

wiley:
11-26-2010

hey there Anders,

would like your opinion on something.

NO2(g) + HCl(g) ---> HNO3(l) + NOCl(g)

http://img225.imageshack.us/img225/6619/hno3.jpg

What you're looking at are 4 simple airtight plastic containers (or 3 if you just vent the NOCl into the atmosphere).
Two containers that produce the gases NO2 and HCl that are directed into a third container(reactor if you will) where the gases react immediately to form HNO3 and NOCl.

What I'm not sure off is whether KNO3 + HCl + Cu will produce NO2 or NO. If NO than it will only produce NO2 until there's no more oxygen in the container and I'd need a different way to produce NO2. Something that actually produces NO2 and no NO that reacts with the surrounding air to make it appear that its producing NO2.

Another thing I'm not sure about is that what if I have an excess of either gas? Will it dissolve in the HNO3? I know NO2 readily dissolves in HNO3 so that leaves HCl, does it dissolve as readily in HNO3 as NO2 does?

Anyways, what do you think?

Anders Hoveland:
Complete Answer to your Idea:

"wiley]hey there Anders,

http://img225.imageshack.us/img225/6619/hno3.jpg "

I am glad you asked me about this. I had a very similar idea some time ago too.

First let me give you my thought. I believe nitric oxide might be able to catalyze the oxidation of anhydrous HCl gas by air. The net reaction would be:
(6)HCl + O2 --> (2)Cl2 + (2)HCl*H2O

Now let me inform you of some known reactions.
A solution of chlorine in water will oxidize NO2 into
HNO3. However, if nitric acid is concentrated enough
(probably over 70%) the reverse reaction will happen!
This is the same reaction for the mix of acids that can dissolve gold, known as "aqua regia", The reaction is:
HNO3 + (3)HCl --> (2)H2O + Cl2 + ONCl
The last product is called "nitrosyl chloride".
This reaction could probably be better written as:
(3)HNO3 + (3)HCl --> (2)HNO3*H2O + Cl2 + ONCl

When I am writing HCl*H2O, this is somewhat like a "hydrate". It really exists as H3O+ Cl-, and the proportion of the pure acid to the water is not a 1:1 ratio, as this simplified structure might suggest.

Your hypothetical reaction:
NO2(g) + HCl(g) ---> HNO3(l) + NOCl(g)
aside from the fact that it does not balance, is almost certainly not possible. NO2 will not oxidize HCl if the solution is dilute; there needs to be strong acidity to dehydrate and cause the reaction. However, concentrated nitric acid is a stronger oxidizer than NO2. The reason the that concentrated nitric acid is an oxidizer is because there is an equilibrium:
(2)HNO3 + H(+) <--> H3O(+) + NO2(+) + NO3(-)
The NO2(+) is called a "nitronium ion". This is why concentrated nitric acid is an oxidizer. This ion prefers to react with water:
NO2(+) + H2O --> HNO3 + H(+)
which is why it only has significant equilibrium formation in very concentrated solutions of nitric acid.
In 10% nitric acid, there is only a very very small equilibrium of nitronium ions. They only start to exist in significant proportions above a 70% concentration.

"What I'm not sure off is whether KNO3 + HCl + Cu will produce NO2 or NO."
While I have not actually tried this, I am almost certain that you could get nitric oxide, and possibly some NO2 if the HCl was 32% concentrated. For NO2 to reduce HCl, the acid must be very concentrated (over 70%), and under normal conditions, HCl will not even dissolve in water over 35%.

"Something that actually produces NO2 and no NO that reacts with the surrounding air to make it appear that its producing NO2."
If it "appears" to produce NO2, it is producing NO2. What other brown gas did you think could be made?
If you only got nitric oxide (NO), this could simply be allowed to react with air, since it spontaneously is oxidized to NO2 by oxygen. The only thing is that you might possibly want to dry the air before it comes in contact with the nitric oxide. As a side note: nitric oxide is colorless and can exist for several minutes before it disproportionates:
3NO --> NO2 + N2O
This means that if you do not quickly oxidize your nitric oxide, you will only be able to get a third as much NO2 from it. I am sending a second letter also.

wiley:
11-27-2010

Hey Anders,

that was quite a lengthy reply. I'll have to read it over a couple more times

Btw, you said:

" NO2 will not oxidize HCl if the solution is dilute"

What solution? The HCl is in gaseous form.
Read this http://www.freepatentsonline.com/4557920.html
They can explain it better than I can

Anders Hoveland:
First, you should not immediately accept every patent as indisputable truth.

Nevertheless, the chemistry in this patent seems quite possible. You should note that EXCESS nitrogen dioxide is used, because if there is extra HCl, it would react with the concentrated nitric acid. I am also somewhat unsure, since I would think concentrated HNO3 would be a stronger oxidizer than NO2, so I would think that the nitric acid would only be of 70-84% concentration, and that there would be some water present.

If you wish to dry HCl gas, I think you can pass it through baked and powdered dry CaCl2, which should pull out any moisture. My local pharmacy store actually sells dry CaCl2, but it is hard to find. If you ask someone that works there, they will have no idea what you are talking about, you will have to thoroughly search every product in the store.

Also, I saved you some trouble, and one of the portions of your proposed reaction. I added some
solid KNO3 to 31% HCl acid, then added copper wire. Nothing happened, but I realized this was due to the KNO3 having difficulty dissolving. I think a protective coating of KCl was forming on the surface of the solid KNO3, and this KCl is not very soluble in HCl because of "the common ion effect".

Next, I added some ammonium nitrate to 31% HCl.
This easily dissolved. Next, I added some copper wire.
First nothing happened, but after a minute tiny bubbles started appearing and the wire started giving off a tiny yellow cloud in the otherwise clear solution.
This reacted for thirty minutes, and the solution turned yellow and a little pile of black debris appeared at the bottom under the wire. The reaction was allowed to continue for 8 hours, and the copper wire still did not fully dissolve, although it was obviously thinner than before. I should also mention that I never saw any brown gas collect over the solution.
When I removed some of the yellow solution and added it to dilute H2O2, the yellow color did not dissappear, indicating the this yellow color was probably not due to dissolved nitrogen dioxide.
I think Cu+1 might have been forming, since Cu+2 is greenish-blue.

So concentrated HCl solution and a nitrate salt do react with copper, but only very slowly. It seems that hydrochloric acid is not concentrated enough for use as a nitrogen dioxide generator.

I also discovered that 30% HCl with 30% H2O2 is capable of rapidly burning copper! It gave off lots of bubbles, and the copper wire was completely gone in less than two minutes, leaving behind a light blue solution.

wiley:
11-29-2010

Hmm... I saw a video on youtube where a guy used a nitrate salt, HCl and copper to produce NO2 gas. He bubbled it through H2O2 to make dilute HNO3. Ah, here it is http://www.youtube.com/watch?v=2yE7v4wkuZU
You can see the NO2 gas in the video.

In the description he says:

"You can use other concentrations of hydrochloric acid but you need to decrease the amount of water added to keep the concentrations the same."

I don't know how critical the concentration is, but maybe that had something to do with why you didn't get much NO2 gas?

Btw, if you know of other ways to generate NO2 gas (that doesn't require HNO3), I'm all ears.

Also, the reason why I'm not distilling HNO3 is because I broke my still a while ago and I'm too broke to buy a new one. I've been making HNO3 the brainfever way and using improvised stills and I don't like neither method. Is the method I proposed to you initially as easy as I think it might be? (provided we can find a good way to generate NO2 gas) As in, 3 airtight polypropylene containers, two of which generate the gases and they're connected to the third container (the reactor) using polyethylene tubing. A very crude setup, but if it were to work (even if it only made 70% conc, acid) it would be an easy way to make HNO3.

Anders Hoveland:

Perhaps my hardware store bought HCl is not really the 31% concentration it claims. Or perhaps using ammonium nitrate instead of potassium nitrate had something to do with it. I did not use very much ammonium nitrate, perhaps I should have used more.
Using KNO3 seemed troublesome, because it did not seem to dissolve in HCl solution, and I did not want to have to add a solution of KNO3, because that would add a big quantity of water, which would dilute the acidity. Solid KNO3 is slow to dissolve in even boiling water.

"if you know of other ways to generate NO2 gas (that doesn't require HNO3),"
Yes, there are several, but most require distilling.
If you can obtain sodium nitrite, mix this with some dissolved nitrate and then add acid. If nitrate is not used, you will get a mix of NO and NO2, and it will require more sodium nitrite, which tends to be more expensive.
NaNO2 + KNO3 + 2HCl -->NaCl + KCl + (2)NO2 + H2O

"Is the method I proposed to you initially as easy as I think it might be? (provided we can find a good way to generate NO2 gas)

Concentrated nitric acid attacks plastic. I do not know how fast 70-80% HNO3 will attack plastic, but 60% concentration can be placed in polyethylene for a short period of time, and 40% can be stored for several months before the acid burns a hole and it leaks out everywhere! I can tell you that my bottle of 70% HNO3 emmitted fumes that gradually burnt through the plastic cap that came with the glass bottle from the chemical company! This took about two years, and consider that the container was stored upright and the acid was never actually in contact with the cap!

Rubber stoppers used when making NO2 gas to make more concentrated acid from only 60% HNO3 acid, were decently corroded (turned into gunky paste on underside) at the end of the experiment.

Also, generating DRY HCl will be somewhat difficult.
For example, using plain Cl2 gas direct from the generator contains too much moisture to make tin tetrachloride (since it reacts with water). I tried this.
I got a few drops of SnCl4, which is a liquid, then it quickly solidified indicating it had reacted with moisture in the gas. Again, I would suggest passing it through baked powdered CaCl2.

The reaction from the patent seems very interesting.
I would like to see pictures if you try it. I would be interested to know how concentrated the HNO3 that forms would be.

I also do not know much about NOCl, whether it reacts with air to form brown NO2, or whether it needs to be heated to do this.

wiley:
11-30-2010

I've distilled HNO3 in polypropylene containers and it held up just fine. Being that I did it in plastic containers, I couldn't heat it much and so it took a long time to complete the distillation, but very little acid decomposes this way so it was highly concentrated. So from personal experience I can tell you that PP and PE are fine for at least a couple of uses.

I think I know where to get Calcium chloride. I think Walmart sells it for use in dehumidifiers.

I already have some NaNO2. It's old though, at least 2 years and I didn't put it in a sealed container, so who knows it could be NaNO3 by now

I really wanna try it and if/when I do I'll be sure to take some pictures.

Anders Hoveland:
12-03-2010

I think the HNO3 with dissolved NOCl will be acceptable to use with most nitrations.
The NOCl will only be moderately soluble, and most of it should come out if the HNO3 is simply allowed to stand for a time. Leaving very concentrated HNO3 in an open container, however, absorbs moisture from the air.

I do not know how concentrated the HNO3 would be from this reaction. I think HCl + NO2 also creates some water, and that the final ratio between water and HNO3 will be an exact proportion, my guess is 80% concentrated HNO3 would be obtained. If any more concentrated, I would think the HNO3 would react with the HCl faster than the NO2.

Anders Hoveland:
12-03-2010

The ammount of water should not matter, but obviously do not use too much or the NO2 will just all dissolve in the water, and few bubbles will come out. But use enough so that all the NaNO2 dissolves.

"I was thinking of generating dry HCl gas by dripping HCl onto CaCl2."
I am almost certain this would not work. CaCl2 does not form a hydrate with water, it simply forms a solution. HCl will probably be as soluble in water whether or not there is any CaCl2. I think the HCl must be dried as a gas. Adding 30% concentrated
H2SO4 to 30% concentrated HCl solution will immediately cause HCl gas to be given off. This is MOSTLY plain HCl, but it will also contain some moisture (water vapor). If you try to dry HCl in its acid solution, there is much more water to absorb.
Drying HCl as a gas should be much easier.
Use a fumehood or do this outside, wear safety goggles and long rubber gloves. Do not breathe in HCl, it is somewhat poisonous. Make sure all the seals on the tubing will not leak the gas! Have a bowl of baking soda dissolved in water ready, preparing for any possible accident.

Another way to make HCl could be to make a chlorine generator with bleach and anything mildly acidic, then burn this with hydrogen gas inside a glass container. If you know of electronics, have two wires and an arc gap to constantly ignite the mixture, keeping it burning, otherwise a small explosition can occur. H2 + Cl2 can also spontaneously explode in the presence of sunlight! The explosion is not very powerful, but will pop out the stopper on a test tube.
Making both the H2 and Cl2 only would require 10% acid of whatever type is convenient.

An idea would be to ignite a small piece of Mg ribbon in the glass container, than begin introducing the Cl2, then a few seconds later, introduce the H2, but have the H2 mixed with CO2. (basically add some bicarbonate to the acid/metal mixture). You would not want to mix plain H2/Cl2- this could be dangerous, there should be another inert gas used in the burning as well, such as CO2 or N2. Perhaps
10%Cl2, 20% H2, and 70%CO2 by volume.
The Mg ribbon would keep burning and keep the gases constantly burning at all times.

But do whatever is easiest and safest

Anders Hoveland:
12-04-2010

very impressive, although I do not understand exactly what everything is in that video.
It seems there is a single hose bubbling gas into a clear liquid. What is that? Is that for the nitrosyl chloride? If this NOCl gas is being bubbled into water, it is likely that the solution is going to give of NO and NO2 gas.

I would encourage you not to give up. This might not be a practical way to make nitric acid, but it certainly is an interesting one. Many people would be interested in that video, if you provided a description of what you were attempting. Must you take it down?

That diagram (on the chemistry site) you refer to seems to be drying HCl gas, not HCl solution. The HCl solution appears to give off vapor which rises up the tube, then descends to be dried by the CaCl2 powder. However, it looks from your video, as if you think solid CaCl2 can dry a liquid solution of HCl. This is doubtful. Try again drying HCl as a gas.

Also a comment, this idea is certainly not basic chemistry, and this reaction is very unusual.

wiley:
12-06-2010

That tube is there to equalize the pressure. Hold a water bottle upside down and empty it half way. Then poke a hole in the bottom, turn it upside down again and empty the other half. It'll drain much faster with the hole in there because it equalizes the pressure.

I used 68g NaNO2 and 80g NH4NO3. 1 mole of each. Is that enough NH4NO3 or should I use more?

That patent says that the gases react immediately. I don't know if that's true, but the second the gases enter the reactor there's definitely something condensing on the walls of the reactor. Of course I suppose it could be the HCl and/or NO2 gas reacting with moisture from the air that's initially in the reactor before it gets purged by the gases.
BUT, I could give the gases more time to react by using a "Y" fitting. Right now the gases enter the reactor separately. But by hooking up the tubes of each gas generator to a "Y" fitting, the gases would react as soon as they come together in the Y fitting.
Yeah, I might try that. But first I still need to make better addition funnels.

Anders Hoveland:
12-06-2010

That site seems somewhat ambiguous about whether it is HCl solution or HCl gas. If you look at the diagram, the tube connecting to the reservoir of HCl solution is on top, suggesting it is the gas that is venting off.

Some possibly important thoughts for you next experiment:
use excess NH4NO3. If there is too much NaNO2, then some nitric oxide (NO) gas will result, this might reduce your concentrated HNO3.

make sure excess NO2 is used. It might be possible that too much HCl will not allow any HNO3, since this can act as a reducing agent.

next time, do not vent off the "NOCl". The NO2+HCl reaction might take a few minutes. If there is a vent, the NO2 and HCl might simply be venting off and then bubbling into the water before they can react!

[Edited on 21-9-2011 by AndersHoveland]

quicksilver - 22-9-2011 at 06:02

I think I am going to experiment with this. I like it more and more.

I have used and (albeit in a very small addition) achieved success clearing up red HNO3 w/ Urea. I realize that superficially it appears counterpoise to purpose but the amounts used are very tiny and the results are significant. I am unsure of the origination of this technique (perhaps Federoff) but it has been used with success and is surprisingly common.

If this could work in a lab setting (achieving an azotrope of 77% from a clean sample of 70%) via simple bubbling, it would be fantastic..

Phantom - 28-9-2011 at 08:33

Here, this will help you> http://www.youtube.com/watch?v=TpWwBxsyJok&list=FLjU9mst...

AndersHoveland - 29-9-2011 at 13:43

Anhydrous (100%) nitric acid could also be prepared by bubbling dry NO2 and ozone gases into a 70% azeotropic concentration of nitric acid. There is another thread in this forum discussing the chemical preparation of ozone. For example, persulfate can be cheaply obtained at pool stores, and if gently heated with nitric acid, the resulting gases will contain a small proportion of ozone, which could be directly bubled into a separate flask of nitric acid.

(2)NO2 + O3 + H2O --> (2)HNO3 + O2

Indeed, superconcentrated ("over 100%") nitric acid can even be made from this method. It is basically HNO3 with a little N2O5 dissolved, and is a stronger nitrating regent. Excess ozone will oxidize away any of the brown NO2 gas that is dissolved, so "white fuming" acid can be obtained, as opposed to "red fuming".

(2)NO2 + O3 --> N2O5 + O2

Ozone can also oxidize SO2 to SO3, so pyrosulfuric acid ("oleum") could be obtained from 90%conc H2SO4 by a similar method.
(2)SO2 + H2O + (2)O3 --> H2S2O7 + (2)O2

[Edited on 29-9-2011 by AndersHoveland]

Steve_hi - 29-9-2011 at 14:41

I made some nitrica acid with out vacuum
2 moles of NH4NO3 and 1 mole H2SO4 plumbing grade.
I used an oil bath of olive oil heated to 180°c.

[img]C:\Users\Steve\Pictures\2011-09-18\upload 1.jpg[/img]

[img]C:\Users\Steve\Pictures\2011-09-18\upload3.jpg[/img]

[img]C:\Users\Steve\Pictures\2011-09-18\upload2.jpg[/img]

The 3rd picture you can see the green where I put a piece of copper wire i had to drop a few drops of water to get it to react.
there was a lot of redish brown smoke and it was bubbling it was quite a stron eaction so I think the concentration was quite high.

quicksilver - 29-9-2011 at 15:04

Catalytic processes of ammonia and platinum wool or gauze [in the production of HNO3] are fascinating. I had often wondered if metals similar to platinum could be used in such catalytic synthesis. (As MMO is used in electrolytic cells such as those used in the production of chlorate.) I attempted to find the answer this this via patents and unfortunately came up with very little. I believe that the MMO coating on a foundation of Ti may present more of a problem, yet I am still very interested as it appears to be an industrial method that may be scaled down IF a substitute for expensive platinum could be found.
If anyone does find some information in the future, I would appreciate you posting catalytic methods here and references.

Steve_hi :
If your condenser has a tilt (downward) toward the collection flask you may see a larger yield. I see some value in a vacuum in the distillation of HNO3, however when I experimented my yields were not appreciably greater until I had a more powerful vacuum apparatus than the unit I used (@ 3pA). There was where the use of Kimax glass became important. I was using Bomex, which was quite thin & subject to both sticking and chipping.

[Edited on 29-9-2011 by quicksilver]

Steve_hi - 29-9-2011 at 15:17

Thanks quicksilver ill ry tilting it downwards on my next batch also thanks for the info on the kimax i intend to buy a 24/40 distilation unit i'll makes sure its kimax. Its quite expensive not knowing what to buy and then end up buying twice.

KemiRockarFett - 29-9-2011 at 15:20

Dry Ca(NO3)2 *4 H20 to Ca(NO3)2 and add H2SO4 , filter out the CaSO4 in a chermic filter. Done !

Another one is to take Ca(NO3)2 *4 H20 and add NaHSO4 and the nitric forms at 40 degrees C when the nitrate starts to melt in its own crystal water. CaSO4, NaNO3 and HNO3 will form.

Nitric acid storage

Steve_hi - 29-9-2011 at 16:51

I was just in my lab making some FeClII and checked in on my nitric acid which I had put in a little glass jar i got at the dollar store which has a rubber seal. The rubber has melted away and the jar is all dirty now I gues the nitric acid ate the rubber away.
I was wondering if the plastic bottle which contained the sulfuric acid would be capable of storing the nitric acid.

bonelesss - 29-9-2011 at 21:22

I`m not 100% sure if HD/PE would be capable of storing Nitric acid,every place where I`ve seen Nitric acid sold they use brown glassbottles, even at low concentration (60%)

If you know any better correcet me.

HDPE containers for Acid

Steve_hi - 30-9-2011 at 01:05

Just found this You are right
http://www.calpaclab.com/pages/chart.html
I guess I'll have to buy one of those little fancy salad dressing bottles you siet on the table that has a glass plug until I can find proper glass bottles. The drug stores don't even have glass bottles here I wonder if American Pharmacies sell glass bottles

[img]C:\Users\Steve\Pictures\Chemical_Reference_Summary-small[1].jpg[/img]

Chemical_Reference_Summary-small[1].jpg - 144kB

jamit - 30-9-2011 at 01:38

If you want to know how to make WFNA, check out unintentional chaos's youtube video on nitric acid, its the best one around.

I can't say from personal experience -- though I have made RFNA -- but you need a vacuum pump or source or its impossible.

Well, good luck!

quicksilver - 30-9-2011 at 06:12

Quote: Originally posted by Steve_hi

Thanks quicksilver ill ry tilting it downwards on my next batch also thanks for the info on the kimax i intend to buy a 24/40 distilation unit i'll makes sure its kimax. Its quite expensive not knowing what to buy and then end up buying twice.

I have distilled my share of acid and you can work with no vacuum with no problem what so ever. When you are dealing with 250ml or less in your collection flask (as projected output) the value is questionable. but when you have a 2L flask and are working for larger yields that percentage becomes very valuable.
Anything but a caustic chemical laboratory-level vacuum will get ruined in about 3-5 distillations because most anything but Teflon-type diaphragm and seals simply won't hold up to the gasses and heat. You could use a water-bed pump HOWEVER - if you don't have a damn good water pressure system it can spit back into your glass-ware and RUIN all your hard work!

If distilling with an alkali metal nitrate and H2SO4 ( like potassium nitrate or whatever) grind down any prills or lumps so the solid nitrate is very finely powdered and mix completely with your H2SO4 till you have a totally clear solution BEFORE hand!!! Then refrigerate so when you start you will be working with a clear solution of mixed acids. That is generally the best method to get every last drop of HNO3 from your efforts. If it's a little yellow: simply add a half gram of Urea to every 500ml of nitric acid and it will clear it up water white.
Dumping powder and your sulfuric into your distillation rig without first having mixed them is a damn good way to plug up your system, get some serious bumping and generally, it's poor lab technique. Get a good clear solution before-hand. You will be working with a standard mixed acid rather than a mess, hoping it doesn't bump or travel into the condenser. A clear solution of mixed acids will also solve a few problems down the line in some nitration than have "mystery" yield or product issues.

The purchase of a simple aquarium air bubblier to push a little dry air through it will work also but you need DRY air. If you keep your collection flask COLD, you will also get a percentage more acid....All the little things are what gets you a good yield. I've worked with stainless steel also in setting up distillation rigs (on a larger scale) and the material is fine if you keep it clean. Since you can't see into any SS distillation rig always flush the whole deal after use with pressure. Welding your joints is the only issue that makes SS distillation a problem and since ACE or ACG sells distillation rigs for $100-150 it's better just to buy them (IMO). I also buy (generally) 24/40. I think that you should always get a tube of "glass grease" or at least make sure you're connections have SOME grease (like a tiny bit of Vaseline up at the top so it's away from the acid). Never connect them dry; even to setup for measurement........they can stick so fast when they're dry; it will break your heart.

[Edited on 30-9-2011 by quicksilver]

Steve_hi - 1-10-2011 at 14:35

I just read this in wicki
Ammonium nitrate gives ammonium chloride and nitric acid upon reaction with hydrochloric acid:
NH4NO3 + HCl → NH4Cl + HNO3

Does this mean i can do a distillation using HCl and get nitric acid as well as ammonium chloride because HCl is much cheaper than sulphuric acid?

VladimirLem - 1-10-2011 at 22:26

Quote: Originally posted by Steve_hi

I just read this in wicki
Ammonium nitrate gives ammonium chloride and nitric acid upon reaction with hydrochloric acid:
NH4NO3 + HCl → NH4Cl + HNO3

Does this mean i can do a distillation using HCl and get nitric acid as well as ammonium chloride because HCl is much cheaper than sulphuric acid?

at the ureanitrate synthesis, this way works, but i dont think, that you can destillate hno3 this way...
The HNO3 would react with the HCL at the round-bottom flask to any other stuff

Just use google...

AndersHoveland - 1-10-2011 at 23:14

Quote: Originally posted by Steve_hi

First, this would not work for one very simple reason. HCl is much more volatile than nitric acid. The HCl would distill out (in the form of an azeotropic mixture).

But secondly, it could be possible that concentrated hydrochloric acid would cause some of the NH4NO3 to decompose after some time, especially when being distilled. The possibility was discussed here:
http://www.sciencemadness.org/talk/viewthread.php?tid=17186

Simply mixing cold 30% HCl and NH4NO3, however, does not lead to any obvious observable reaction.

[Edited on 2-10-2011 by AndersHoveland]

Luftwaffe - 1-10-2011 at 23:28

I do not believe distilling nitric using HCL would be possible considering Aqua Regia is the mixture of HCL and Nitric Acid.

Consider the reaction,

HNO3 (aq) + 3 HCl (aq) → NOCl (g) + Cl2 (g) + 2 H2O (l)

As well as,

2 NOCl (g) → 2 NO (g) + Cl2 (g)

I believe your yields would be bad and would require a very good scrubber or ventilation to deal with all the toxic and corrosive gasses. Even if you tried to distill nitric from a mixture of HCL and a nitrate salt it would take so much time to get a reasonable amount of product. Time is money, and I think spending a few extra bucks on some drain cleaner is better than a few days of frustration.

Sources: Wikipedia, personal knowledge.

AndersHoveland - 2-10-2011 at 01:22

Quote: Originally posted by Luftwaffe

I do not believe distilling nitric using HCL would be possible considering Aqua Regia is the mixture of HCL and Nitric Acid.

HNO3 (aq) + 3 HCl (aq) → NOCl (g) + Cl2 (g) + 2 H2O (l)

Actually, that reaction is an equilibrium reaction. It tends to go to the right when the acids are very concentrated, but more often the reverse reaction happens when the reactants are more dilute. NOCl actually hydrolyses with water to nitrous acid and hydrochloric.

NOCl + H2O --> HNO2 + HCl

In the presence of water, chlorine easily oxidizes nitrous acid.

HNO2 + H2O + Cl2 --> HNO3 + (2)HCl

So the net reaction is just the reverse, assuming of course that the acids are dilute. Even in 30% concentrated acids, the equilibrium primarily exists in the form of HNO3 and HCl, although there is enough equilibrium to attack gold.

All too often, seamingly simple reactions are actually much more complicated. Most of the time, for convenience, we like talking about the simple net reaction, but these types of generalisations do not hold true for all conditions.

Steve_hi - 2-10-2011 at 09:57

Just checked in on my nitrica acid today Ive been keeping it in a small fridge that wasnt turned on when I opened the door I stunk the whole place up. The nitric acid ate all the rubber seal and turned it to mush and was eating the steel bail that keeps the lid down. I think the first thing to do in the synthesis of Nitric acid is to have a suitable container to store it in.

[img]C:\Users\Steve\Pictures\2011-10-02\1.jpg[/img]

quicksilver - 2-10-2011 at 11:55

Teflon (type) seals / stoppers or "glass to glass" (roughened glass stoppers)..... Few other materials will appropriately seal HNO3 for any period of time. But keeping it cool and away from UV (dark amber glass) IS a smart idea.

Steve_hi - 2-10-2011 at 15:30

I had an idea this afternoon I was thinking to use a brown beer bottle and machine a flat piece of aluminum the size of the OD of the beer bottle neck and put it inside the bottle cap since Aluminum is not reacted by nitric acid it should suffice as a seal ?? Not that I dont want to buy a proper bottle but pro lab scientific wants 15 dollars for a bottle and it is only hdpe cap plus the shipping

AndersHoveland - 2-10-2011 at 21:06

Quote: Originally posted by Steve_hi

... since Aluminum is not reacted by nitric acid it should suffice as a seal ??

Aluminum is actually resistant to nitric acid if the acid concentration is over 95%, but if the concentration is below 80%, or if the nitric acid is heated above 40degC, the corrosion rate is much faster.
This is only for handling nitric acid. If the nitric acid is to be stored for long lengths of time, aluminum is probably not advisable, as it no doubt would gradually corrode over time.

Luftwaffe - 2-10-2011 at 21:45

Instead of aluminum, just line the cap with teflon tape. I intend on storing my nitric with glass perfume bottles with teflon tape lining the plastic cap. I buy my reagents from prolab, I was quite dissapointed over there reagent bottles so I have to improvise.

AndersHoveland - 3-10-2011 at 00:07

would like to share an interesting reaction from the literature of Marcellin Berthelot, dry ammonia gas reacts with the nitrogen dioxide and nitric oxide, at room temperature,

(2)NO2 + (2)NO + (4)NH3 --> (2)NH4NO2 + (2)H2O + (2)N2

Solid ammonium nitrite inside a tube explodes if heated on a water bath to between 60-70degC. And the substance gradually decomposes at room temperature, slower if cold, or faster in aqueous solutions, forming nitrogen gas.

Nitric acid from CAN

Steve_hi - 6-10-2011 at 12:27

Someone mentioned that CAN "Calcium Ammonia nitrate" could not be or would be difficult to use as the nitrate for this synthisis but to day i repeated my previous distillation with NH4NO3 that I got from CAN fertilizer and I had no problem with it, it worked fine.

[img]C:\Users\Steve\Pictures\2011-10-06\1.jpg[/img]

hissingnoise - 7-10-2011 at 01:11

Hot, concentrated sulphuric acid is a good oxidiser and will oxidise the ammonium ion, introducing water and diluting the distillate!
Use an alkali-metal nitrate for the highest concentration of HNO3 . . .

Steve_hi - 7-10-2011 at 01:25

I have a number of different nitrates can you mix the nitrates and do this synthesis. like KNO3 and CaNO3 and I have BaNO3 I ordered by mistake just curious bout all these diffent nitrates I have and dont know what I will do with them

Got overexxcited about learning chemistry when I was ordering chemicals

hissingnoise - 7-10-2011 at 03:50

The sulphates of barium and calcium being insoluble will complicate things.
K and Na nitrates are ideally suited for preparing HNO3.

Luftwaffe - 7-10-2011 at 04:24

If you use CAN fertilizer, make sure you know the percentages of each nitrate or you may be using too much of an excess or too little sulfuric acid in your nitric synthesis. Basic stoichiometry I know, just making sure!

quicksilver - 7-10-2011 at 06:14

hissingnoise :

Some interesting idiosyncratic elements about CAN. The Yara product may have been one of the few ammonium nitrate products to halt the use of a FGAN to be used as an ANFO or energetic. A great deal of research was put into the final Norwegian export product. When exposed to water the Yara product yields a "Glue like material" that clogs a filter. It's a Carboxylate polymer [and] is somewhat a generic term as it could mean China Clay and oil or a synthetic. It is obviously a “trade secret” as it’s not totally spelled out in their patent. The original is Norwegian and I couldn't get an English copy. The US was 6,930,139 and gave enough info to develop two mechanical methods of extraction. And there may be a chemical method of extraction but the recovered material would be KNO3 rather than the original nitrates. The mechanical processes work well but one is somewhat slow and the other requires a moderate amount of tools in construction. I had some fun figuring out workable extraction techniques: however I take my hat off the Yara as their prilled product (IMO) would be safe for sale unless the user to some real time in examining the techniques to isolate the original NH4NO2.

From a utilitarian stand-point, Yara may have stopped the common abuse of NH4NO3 from dangerous, unskilled use as a blasting agent. The market they have at this time is very large. NH4NO3 will only be sold [in configuration similar] as a bagged product. It is still available to agriculture via truck (unbagged- delivery) but there will be a minor registration process - all laws available via the DHS page, etc.
This IS a multi billion dollar market, when all is said and done.

The interesting thing was that many people have attempted to put the breaks on bagged ammonium nitrate access by the public and the Yara company may have been the first to come close. They have succeeded in grabbing the US/CA market of bagged FGAN, which is quite an accomplishment.
The Dept. of Homeland Security (DHS) has proposed ('08) & got enacted into law a national focus on NH4NO3. If you ever do want to check out what lengths the DHS had gone to so as to grab a headline or two and give themselves a little promotion check out: "the Secure Handling of Ammonium nitrate Program" DHS, (see attached summery), the bulk can be Googled as well as the specifics.

The new laws for the US deal with bagged NH4NO3 and are already in effect. but as many people have said through out the years it would be consistently possible for a skilled / clever individual to be able to extract ammonium nitrate from a prill-based starting material. It's just that Yara had come the closest so far.

Attachment: AN_ANPRM.pdf (254kB)
This file has been downloaded 785 times

[Edited on 7-10-2011 by quicksilver]

Steve_hi - 7-10-2011 at 07:23

Interesting how goverment agencies and the general public can be so stupid.
How many people have been blown up by NH4NO3
How many people have died from tobacco
How many people have died from hand guns
How many people have died from alcohol and families ruined in the process
How many people have died from patent protection barring people from being able to afford life saving medications , simply to protect the wealth of the richest people in the world.\

Trudeau said "The public is an ass" HE WAS RIGHT

hissingnoise - 7-10-2011 at 08:08

Quote:

How many people have been blown up by NH4NO3

It's AN's potential for blowing people up rather than the numbers but historically AN does have a reputation.
State agencies can attempt making its extraction 'difficult' but that's all they can do!
Small hurdles for anyway committed terrorists!
And they know it . . .

quicksilver - 7-10-2011 at 11:11

Quote: Originally posted by Steve_hi

I had an idea this afternoon I was thinking to use a brown beer bottle and machine a flat piece of aluminum the size of the OD of the beer bottle neck and put it inside the bottle cap since Aluminum is not reacted by nitric acid it should suffice as a seal ?? Not that I dont want to buy a proper bottle but pro lab scientific wants 15 dollars for a bottle and it is only hdpe cap plus the shipping

Teflon stoppers ARE available and are worth the effort to buy. As you know HNO3 absorbs water fairly quickly from almost any exposure. While amber glass is a good idea, I would do all I could to achieve a safe and air-tight seal. What's more you certainly don't want them to break & thin commercial glass may not be the best choice; more so because they are designed poorly and tip over easily.
Common "black rubber" stoppers simply can't be sealed effectively against a strong acid's fumes and the mess they make when decomposing is disgusting.
There are polyurethane stoppers that have been wrapped in Teflon tape and appear to last for a few weeks but they also decompose. Teflon stoppers are really the best choice. They aren't cheap but they will last for many years of active use in most common caustics.

magnus454 - 7-10-2011 at 20:08

Remember this is bad stuff, they mix it wife jet propulsion kerosene to make ICBM rockets fly! My father dealt with this stuff in the US ARMY guided missile division.

Eyro9001 - 7-10-2011 at 21:21

I have yesterday manufacture of nitric acid by distillation way
The quantity = 17ml
weight = 27g
Density = 1.59
!!!!
(the color is very yellow)
How much is the concentration of the acid?
Is there a problem?
How do I know the presence of oxides of nitrogen in acid or not

Note
This my first experience in the manufacture of nitric acid

Thanx >>>

[Edited on 8-10-2011 by Eyro9001]

Luftwaffe - 7-10-2011 at 22:07

I shall give you the answer, UTFSF! Concentration can easily be found from density charts on the internet and oxides of nitrogen can be seen by mere visual observation. You can even use NOx scrubbers to rid of all traces of them which has been discussed in many posts on SM, search to ascertain the answers you seek.

Eyro9001 - 8-10-2011 at 03:00

Luftwaffe :

Quote:

Concentration can easily be found from density charts on the internet

Yeah I know that, but the density is 1.59 ! not found in table of concentrations. (that mean up of 100%!!!)

Quote:

oxides of nitrogen can be seen by mere visual observation

Explain this point please, or give me the links
Thank you very much

Steve_hi - 8-10-2011 at 06:10

I think that the fact it is yellow in colour signifies NO because pure HNO3 is clear in colour

hissingnoise - 8-10-2011 at 06:56

Gaseous NO2 is a dark reddish brown colour but the dimer N2O4, existing at low temperatures, is colourless . . .

Luftwaffe - 8-10-2011 at 14:30

Eyro, how precise and what volume of acid were you measuring? When calculating densities even a small inaccuracy is a big deal if your only measuring a few mL of acid. Perhaps scale up your measurement and see if your results change.

As for the equilibrium between NO2 and N2O4, you can shift the equilibrium towards making NO2 by chemically removing NO2 with a scrubber forcing the reaction to completion.

N204 --> 2 NO2

[Edited on 8-10-2011 by Luftwaffe]

hissingnoise - 9-10-2011 at 05:46

In this case, the equilibrium is solely temperature-dependent.

Kiwichemicali - 25-3-2012 at 00:52

Quote: Originally posted by quicksilver

If it's a little yellow: simply add a half gram of Urea to every 500ml of nitric acid and it will clear it up water white.

[Edited on 30-9-2011 by quicksilver]

I've distilled Nitric Acid from Sulfuric Acid and Potassium Nitrate. The result is light yellow. However when I add Urea a vigorous reaction takes place en lots of NOx fumes escape.
Is the addition of Urea to fuming Nitric Acid always this violent?

Kiwi

Pulverulescent - 25-3-2012 at 02:57

ISTR some fizzing but no NO2 evolution.
Urea decomposes nitrous acid in solution to N2 and water, diluting the acid.
Perhaps you used too much of it . . .
A better solution is bubbling dry air through your HNO3 to oxidise NO2.

AndersHoveland - 25-3-2012 at 10:00

This is why I do not think urea is a good way to remove NO2 from nitric acid.

The most ideal way would be to pass dry air through an ozone generator. Then the ozone could oxidize the NO2 to nitric acid, increasing the concentration even further.
(the reaction of plain oxygen with NO2 only forms HNO3 up to a certain concentration, whereas O3 can even oxidize NO2 to N2O5, the acid anhydride of nitric acid.

[Edited on 25-3-2012 by AndersHoveland]

Kiwichemicali - 26-3-2012 at 08:18

Quote: Originally posted by AndersHoveland

This is why I do not think urea is a good way to remove NO2 from nitric acid.

The most ideal way would be to pass dry air through an ozone generator. Then the ozone could oxidize the NO2 to nitric acid, increasing the concentration even further.
(the reaction of plain oxygen with NO2 only forms HNO3 up to a certain concentration, whereas O3 can even oxidize NO2 to N2O5, the acid anhydride of nitric acid.

[Edited on 25-3-2012 by AndersHoveland]

Any idea how much air needs to pass through 100ml? I tried pure O2 but gave up after a short time because I didn't see a result. I really have no idea what volume or how much time is needed.

Kiwi

Pulverulescent - 26-3-2012 at 08:25

It's slow, takes patience ─ and warming the acid helps . . .

AndersHoveland - 27-3-2012 at 01:06

Quote: Originally posted by Kiwichemicali

Any idea how much air needs to pass through 100ml? I tried pure O2 but gave up after a short time because I didn't see a result. I really have no idea what volume or how much time is needed.

1 liter of gas contains 0.0446 moles of molecules under normal conditions. 100ml liters of water contains 5.556 moles. In the reaction of NO2, O2, and H2O, one mole of O2 makes the equivalent of 4 moles of HNO3, or when further concentrating the HNO3, 1 mole of O3 makes 2 moles of HNO3. You can do the calculations. A rough estimate is that 100 liters of gas contains the same number of molecules as 100ml of water. So that is a large quantity of gas, and the reaction will take long.

Vikascoder - 27-3-2012 at 21:31

Quote: Originally posted by KemiRockarFett

Dry Ca(NO3)2 *4 H20 to Ca(NO3)2 and add H2SO4 , filter out the CaSO4 in a chermic filter. Done !

Another one is to take Ca(NO3)2 *4 H20 and add NaHSO4 and the nitric forms at 40 degrees C when the nitrate starts to melt in its own crystal water. CaSO4, NaNO3 and HNO3 will form.

what will be the concentration of HNO3 formed after filtering

hames - 17-4-2012 at 23:55

just a quick question if anyone can help,I was wondering if you can use 85% phosphoric acid to make nitric acid from a nitrate salt rather than sulphuric acid,I have both but my sulphuric is expensive and I get litres of the phosphoric for free and want to get some good use out of it.

weiming1998 - 18-4-2012 at 00:50

Quote: Originally posted by hames

just a quick question if anyone can help,I was wondering if you can use 85% phosphoric acid to make nitric acid from a nitrate salt rather than sulphuric acid,I have both but my sulphuric is expensive and I get litres of the phosphoric for free and want to get some good use out of it.

Ca(NO3)2/other metal nitrates that forms insoluble/slightly soluble phosphates might work, but KNO3, NaNO3 and NH4NO3 definitely doesn't work. Try it with a bit of calcium nitrate and a bit of phosphoric acid first.

[Edited on 18-4-2012 by weiming1998]

hames - 18-4-2012 at 01:45

Quote: Originally posted by weiming1998

Quote: Originally posted by hames

I would be distilling it afterwards so it shouldn't matter if the salts are soluble should it, I have an abundance of KNO3,I have no Ca(NO3)2 but I have Ba,Sr nitrates which might form insoluble phosphates.

AJKOER - 21-4-2012 at 05:47

I was wondering whether anyone has prepared concentrated HNO3 from the action of Oxalic acid on a nitrate salt:

2 KNO3 + H2C2O4 ---> K2C2O4 + 2HNO3

Potassium oxalate may be soluble in dilute Nitric acid but insoluble in concentrated or in cold concentrated solutions. For feasibility per "A dictionary of chemistry and the allied branches of other sciences", Volume 4, by Henry Watts, page 251, to quote. Oxalate acid "is one of the strongest acids, decomposing dry chloride of sodium when heated, with evolution of hydrochloric acid, and converting chloride or nitrate of sodium in aqueous solution into acid oxalate.
The oxalates of the alkali-metals are soluble in water, the rest are for the most part insoluble in water, but soluble in dilute acids. "

Link:
http://books.google.com/books?pg=PA255&lpg=PA250&dq=...

An indirect approach may also be successful. Add copper turnings to a heated concentrated solution of H2C2O4 and KNO3. Collect the NOx gases, mix with air and dissolve in water.

If dilute HNO3 is presence:
3Cu(s) + 8HNO3(aq) ——> 3Cu(NO3)2(aq) + 2NO(g) + 4H2O(l)

If Conc HNO3:
Cu(s) + 4HNO3(aq) ——> Cu(NO3)2(aq) + 2NO2(g) + 2H2O(l)

[Edited on 21-4-2012 by AJKOER]

quicksilver - 21-4-2012 at 08:42

Oxalic acid isn't as "thirsty" as sulfuric acid. It's possible that some of the water could be retained but not as powerfully. What's more it's a powder so if you used it with a nitrate you would have to use water (from the HNO3 ?). Therefore you may have a 3rd adulterant chemical in a mixed acid wherein you simply wanted it for a clean nitration (or whatever). But if you tried it with HNO3; I don't believe you could do better than cooked and cleaned sulfuric (@ 98%), However it's a good question because to think along those lines MAY discover an acid that would preform better (than H2SO4). This is a very interesting "micro" experiment where weights and volumes after separation may be become a "hope note" that continues for quite some rime.

hames - 22-4-2012 at 23:31

I know phosphoric acid is used to make hydroiodic and hydrobromic acids because it doesn't oxidize them to their elements like sulphuric acid does.is there any reason this won't work for nitric acid? you would think at least the azeotrope would be formed.

AndersHoveland - 25-4-2012 at 17:11

Nearly anhydrous nitric acid apparently can be made by carefully distilling nitric acid with dry phosphoric acid. It is important not to add the phosphoric acid all at once, as the heat can cause decomposition of some of the nitric acid. (Weber)
http://books.google.com/books?id=CrNXAAAAYAAJ&pg=PA342&a...

AndersHoveland - 19-3-2013 at 15:03

For any of you thinking about dissolving N2O5 into nitric acid to raise the acid concentration, this may be of interest:

Nitric acid dissolves nitrogen pentoxide, and a definite compound, 2HNO3.N2O5, has been obtained which is liquid at ordinary temperatures but solidifies at 5° C.

caterpillar - 21-3-2013 at 13:55

If you need WFNA for RDX production, you can successfully do without it. Distillation of KNO3 (NaNO3) + H2SO4 gives you yellow liquid, which most likely will burn hexamine, producing red fume of NO2 (I got it during my first experiments). But there is the simple way to resolve this problem. Dissolve NH4NO3 in HNO3 and use hexamine dinitrate instead of pure hexamine. Two birds with one stone- forget about nitrogen oxides and increase total yield. Read topic on RDX.

Ral123 - 21-3-2013 at 23:31

I've never heard adding NH4NO3 helping increase yields of RDX toward WFNA. Why is nobody doing it?

caterpillar - 22-3-2013 at 01:23

"There are more things, on the Earth, on the Heaven, Horatio..."

Organic Chemistry of Explosives by Jai Prakash Agrawal & Robert Hodgson.

This method is known as the K-process after its discoverer Kofﬂer.

Like method 5.15.1.2 it
uses ammonium nitrate to compensate for the nitrogen deﬁciency in hexamine and works to
Equation (5.24) where two moles of RDX are produced per mole of hexamine. As observed
with method 5.15.1.2, the addition of ammonium nitrate to nitric acid appears to prevent
dangerous oxidation reactions from occurring. In fact, this nitrolysis reaction only occurs at
elevated temperature and so a constant temperature of 80 ◦
C is usually maintained throughout
the reaction. Yields of approximately 90% are attainable based on one mole of hexamine
producing two moles of RDX.

I can only add that if you use hexamine dinitrate, you save some cons HNO3 because part of acid is added as a part of hexamine dinitrate. And even without NH4NO3 using dinitrate prevents dangerous oxidization

AndersHoveland - 24-3-2013 at 00:18

Quote: Originally posted by caterpillar

As observed
with method 5.15.1.2, the addition of ammonium nitrate to nitric acid appears to prevent dangerous oxidation reactions from occurring.

Perhaps because the ammonium ions help absorb any undesirable NO2 that forms. Normally this is a slow reaction and requires boiling temperatures, however, so I am not completely sure.

Ral123 - 24-3-2013 at 00:44

Ammonium nitrite is so unstable, ammonium chloride+sodium nitride yields nitrogen. I also have the opinion that with HDN the reaction goes more smoothly. I don't need to economise the hexamine witch is sold by the kilo quite cheap. The expensive here is the WFNA an all the solvents and distilled water for the HDN and washing and purifying the RDX.

caterpillar - 24-3-2013 at 02:12

Well, that's just the same situation what I had, trying to prepare RDX. Conc HNO3 is the most valuable component here. This is why any way to replace some of HNO3 with diluted one (or even mixture of H2SO4 + KNO3 (NaNO3) is more than welcome. And this is the way to do it- one need DILUTED HNO3 to prepare hexamine dinitrate. Therefore, part of conc HNO3 can be replaced- it goes into reaction vessel as part of hexamine salt. Of course, this salt must be dry. Well, large excess of HNO3 is a must, but later diluted HNO3 may be used for preparation of hexamine dinitrate. On the other hand, I have such feeling, that using this method one can put more hexamine per a gram of HNO3 than usually (only hexamine and HNO3). And there is another way- Anders gave me link to it. A way, which lets one to use mixed acid instead of pure HNO3. I only suspect, that this method requires oleum, not only conc H2SO4. About which process is smoother. My own experience- reaction with hexamine dinitrate is milder out of question. Later, if you need to get more pure and stable RDX, you should perform oxidation of all linear nitramines. You must use reaction vessel, which can withstand fast heating (like pirex). The most civilized way to perform this op is to put slowly reaction mixture (70%) and water(30%) into same vessel. What you'll get is the typical runaway- mixture will boil. Do it in open air- NO2 will go away. But finally you'll get large crystals of RDX. Wash them with water.

greenlight - 25-2-2015 at 07:10

Does anyone know if one of these Dreschel bottles can be used like a cold trap to neutralize acid vapours when distilling Nitric acid under vacuum.
The one in the picture is mine and I am unsure if it can withstand the pressure of my Diaphragm pump? Has anyone used one for this purpose?

[Edited on 25-2-2015 by greenlight]

[Edited on 25-2-2015 by greenlight]

Hawkguy - 25-2-2015 at 09:02

Quote: Originally posted by hissingnoise

If you need water-white HNO3 just add a small pinch of urea as a last resort.

Won't the resulting Urea nitration cause the release of more NOx, which might be redissolved?

Molecular Manipulations - 25-2-2015 at 09:38

greenlight, I have used those before under about 50 torr, it didn't implode and was NOT vacuum rated. Really though nobody knows your glass better than yourself, so don't take my word for it. I suggest testing it under vacuum before you try the distillation. Losing that piece would suck, but it'd be better than losing it along with a bench or getting an acid burn if it implodes with fuming nitric acid in it.
How much vacuum can your diaphragm pull?
Are you going to neutralize the vapors with a basic solution or just condense them? I suggest the latter, but that's just because I use a real vacuum pump which is very expensive and could be ruined by acid vapors. A cold trap won't do a whole lot for WFNA at 50 torr.

markx - 25-2-2015 at 12:17

For low pressure distillation of corrosive substances like WFNA I would warmly reccommend the usage of a water jet vacuum pump. Preferably configured in to a pump driven recirculation system that drives the same water through the venturi jet in a loop. No need to dump hundreds of liters of water down the drain driving a jet pump off the tap directly. The scrubber systems will not save a mechanical vacuum pump from the harm of the corrosive acid or solvent vapours....at best they delay the inevitable death of the device. It is just a matter of time before the scrubbers saturate, the cooling of the vacuum trap fails or for some reason a huge amount of corrosive vapours is sucked through the system too quickly. Either way the harm will enter the pump and shut down the operation in a very uncomfortably short period of time. Unless you have a fully teflon coated pump system at your disposal, but these are rare and quite expensive. Hence a venturi type device is your best option, cheap, simple and quite insensitive towards most kind of corrosive environments.

Besides that the vacuum distillation for WFNA in amateur setup is a rather unnessesary complication and the benefits do not outweigh the added technical complications and dangers. There are very few purposes that I can think of for which a properly performed atmospheric pressure distillation will not yield a suitably pure WFNA (and heterocyclic nitramines are not among those purposes). Hence I would suggest to opt for the least technical complications and focus more on tuning the operating parameters of the system and carefully dividing the received fractions.

[Edited on 25-2-2015 by markx]

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