Sciencemadness Discussion Board

Nitric acid

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a123x - 12-1-2003 at 10:20

I've found a new way to make nitric acidof high concentrations and purity. It requires sulfuric acid, and either Ca, Sr, or Ba nitrate. Its simple really, Ca, Sr, or Ba sulate are insoluble so adding Ca(NO3)2 to sulfuric acid of an concentration and you can get that same concentratio of nitric acid back with CaSO4 precipitating. Then its just a matter of filtering. So if you can get 98% sulfuric you can make 98% nitric with nothing dissolved in it and no distillation required.

TheBear - 12-1-2003 at 10:37

I've also read some about this method. Does anyone have any experience with this method using Ca(NO3)2?

DeusExMachina - 12-1-2003 at 12:19

shouldn't this be in the chemistry forum?

The "precipitation-nitric" - A tale of hope

BASF - 12-1-2003 at 12:21

This method is far from being new.....but i also once gave it a try.
It seemed like a geniously simple method to me at first.......

I´ve experimented a lot with these precipitation-reactions for nitric acid, but experienced that this method sucks heavily.
The only way i could imagine working would be to use a porous glass-filter with vacuum sucking device.
If u experiment a bit yourself you will see very soon that the volume of the developed HNO3 is not enough to decant it off, even with altering the precipitate you will not be able to decant more than a few percent of the whole liquid in the slug.

The reaction product is a thick white paste and it is very important to use fine powdered, dry Ca(NO3)2.
As you mentioned, Ba(NO3)2 is also an option, with slight advance to the other method, because of the extremely low solubility of BaSO4, even in concentrated mineral acids.
-But on the other hand: Ba(NO3)2 is very toxic and expensive, although i have to admit it can be made via:

BaSO4 + 10-fold excess of Na2CO3 with little water; cooking for 1/2hour > BaCO3(insoluble in cold water)

BaCO3 + dilute HNO3 > Ba(NO3)2

Other methods i have been trying were extraction with solvents such as acetic acid and the like.
I had to experience that these multiple extraction methods have a huge problem in the 2nd step when u try to seperate the acetic acid from the nitric acid.

As it was fairly cheap, i first tried extraction of the acetic acid with toluene, which worked fine with 69% commercial nitric acid in the mix with glacial acetic acid.
Later then i had to come to the conclusion that these extraction methods are very very sensitive to contaminations of inorganic salts, such as rests of unreacted Ca(NO3)2 and also unreacted sulfuric acid.

Even the slightest contamination with these, may it be in the order of only a few mg per 100ml cause the whole thing to fail and it won´t mix any more.

But, it has to be mentioned that the E&W-forum discussed some other, though more extensive extraction method using methylene chloride.
Several ways of removing the methylene chloride from the mix can be used, from freezing the nitric out to letting the methylen chloride evaporate.
As long as you can get hold of bigger quantities of methylene chloride it would be fine.

My last thought of how to use the precipitation method was to mix the thick slug with 1 or 2 fold its volume of glacial acetic acid and use it for different nitrations as is.
Or, if u mix it with even bigger volumes of acetic acid, u could yet decant the mix carefully, giving a dilute solution of anhydrous nitric acid in glacial acetic acid which had some advantages to the slug due to being it a clear liquid, not a thin slug........but i fear this would slow down the rate of reaction for nitration and lower the yield.

But at last, this is one of my experimental favourites, and i´m sure some further experimentation on it would show some interesting uses for it...


Polverone - 12-1-2003 at 13:22

One other way of making Ba(NO3)2 is to heat BaCO3 with NH4NO3 and a minimal quantity of water at high temperatures. This method is not very fast, but it doesn't need any nitric acid.

A (perhaps silly) question: can the precipitates be aged to obtain larger crystals? I know that sometimes letting precipitates stand for a long time, especially with mild heating, will cause larger crystals to form instead of the extremely fine particles that you may have initially. I would especially wonder about using this method with calcium nitrate since calcium sulfate is at least somewhat soluble. I wonder if a few days just sitting at (say) 40 degrees Celsius would lead to a larger, more easily filterable precipitate, or even one that liquid might be decanted from.

BASF - 12-1-2003 at 14:15

No silly question at all.....:)

I tried altering the precipitate in the hope(as you mentioned) it could then settle to a higher degree and make decantation of at least some ml of nitric per batch possible.....

So far, this has definitely failed.
It seems to me that the CaSO4-precipitate alters very quickly(i guess its a matter of hours), but the effect on settling was hardly noticeable at all.

But when it comes to filtering, it might be a good idea......i believe it would also extend the life-expectancy of precious glass-filters...

Maybe glass-filters could be cleaned with hot conc. H2SO4 sucked through as CaSO4 is quite good soluble in it making a complex-anion.

Anybody interested in the Ca(NO3)2/H2SO4-method may also take a look at
the Ca(NO3)2-synth in the library, which makes use of NH4NO3-fertilizer/CaO(lime).

a123x - 12-1-2003 at 15:34

I don't have any calcium nitrate, I'm intending to make it with NH4NO3 + CaOH. Not having calcum nitrate I decided to do a test with CaCl2. I made a solution of CaCl2 and added some sulfuric to it which of course resulted in plenty of white precipitate. As a filter I decided to use a model rocket parachute that I had which is made from nylon if I remember correctly. It worked beutifly as a filter, there was actually a large excess of CaSO4 in it which caused it to filter very slowly but just taking a plastic cup and pressing down on the glob of moist CaSO4 squeezed the HCl solution out and through the filter. The filter wasn't damaged or anything, and there was only a very slight amount of CaSO4 that got through. I'm not sure why this wouldn't work to filter CaSO4 from nitric acid unless nitric acid damages nylon.

lucifer - 12-1-2003 at 16:01

Can 100% HNO3 be stored for a long period if kept cool and in the dark or will it very slowly decompose into NO and H20.

BASF - 13-1-2003 at 07:23

lucifer: it won´t decompose under these conditions.

a123x: hmm your method is quite interesting.......maybe one could use a foil of a more resistant material, poly propene("poly propylene, pp") for instance or pvc which would be even more resistant in my opinion, which has little holes in it made by just stinging needles into it(very fine needles could be made of glass pipes).

This reminds me a lot of kneeding dough and cooking after all, hehe

a123x - 13-1-2003 at 12:09

Yeah, as it turns out the HCl mildly damaged the parachute but sulfuric puts holes through it in a few seconds. I don't actually have any nitric handy to test it. The sulfuric is stored in a polyethlyne bottle so I figure that it should hold up to nitric acid too, and thin polyethlene sheets are easy to get. Really a cloth made of those materials would be better though since it would work much more easily as a filter but I can't think of any where that polypropylene, polyethlene, or teflon cloth could be gotten.

BASF - 13-1-2003 at 13:56

i tend to say that PE is not applicable for this action, sulfuric acid is one thing, 100% nitric acid another.....

But why not try PVC?
-It is far from being ideal, but maybe good enough for a single filtering.

a123x - 13-1-2003 at 14:15

Actually I was just looking at a chart about chemical resistance and it seems that PE has a better resistance to 70% nitric acid than PVC does. I might go a different route anyway. I'm thinking of taking some glass and breaking it up into small bits and then makinga filter using that. From what I can tell of the nature of the precipitate/liquid mix that forms it might work. When it isn't too thick the liquid should be able to flow through the filter leaving most of the sulfate at the top. After that point the mix is pretty thick and requires being pressed t squeeze out the liquid where the thick sulfate mixture isn't capable of flowing through the filter. There has to be a way to solve this problem other than distillation.

BASF - 13-1-2003 at 17:39

Strange. I expected PE to be less resistant than PVC to nitric acid but if the chart says so.....

BTW, in the past i have also thought a lot of how to press out the few liquid that the slug cannot imagine what i have tried.....beginning with syringes, paraffine-filled u-pipes, silicone-squeezes.....never ended up with more than a few drops though.

Maybe u could press it better through your "foil-filter" if u wrap it with the foil like a bonbon.....know what i mean?
-This could be a simple method with the advantage of having the slug protected from moisture and also, there would be less poisonous HNO3-fumes to struggle with.

Microtek - 14-1-2003 at 01:59

How about using a centrifuge to separate the CaSO4 from the HNO3 ? There would be some work involved in building the machine, but a centrifuge can be a quite useful thing in a lab... At high enough acceleration, the precipitate is compacted into a pellet at the bottom of the flask.

BASF - 14-1-2003 at 05:09

I guess this could be made with technic-lego(we all have played with it..) which i believe is also a good thing when you want to build your own stirring device or even your own remote robo-hand for handling dangerous substances.....:D

Of course, the plastic parts have to be protected against heat above 50°C, organic solvents, nitric acid and the like, but maybe a coating with wax or some other coating that could maybe be sprayed on it would do part of the job.

Wrapping it in aluminum foil could also help, but does not protect against corrosive vapours, such as those from nitric acid.

Back to the centrifuge.....
Good idea as you won´t need any filtering but i also think the problems are the vibrations, even with well-working counter weight.

Anyone having done some analytical lab course on university or at school, knows what i am talking about - Even the industrial products suffer from heavy vibrations, often moving around the table while working.

a123x - 14-1-2003 at 12:20

A centrifuge would be nice but I don't have access to one and I don't think I'd be very capable of building one. I did finaly think of something easy enough to get that might make a decent acid proof filter, fiberglass insulation. I looked up some information about it online and the only problem would be the phenol-formaldehyde that binds the glass fibers together. I figure if the acid harms phenol-formaldehyde then sulfuric acid might be usable to remove it from the glass, if it is neither harmed by nor harms the acid it could just be left on. The only problem is if it messes up the acid in which case it would need to be removed by solvent. Does any one know if there is any particular solvent that would be best at dissolving phenol-formaldehyde?

BASF - 15-1-2003 at 09:55


Does any one know if there is any particular solvent that would be best at dissolving phenol-formaldehyde?

I have no specific reference for this, but i guess phenol-formaldehyde could eventually be dissolved in glacial acetic acid.

DeusExMachina - 18-1-2003 at 17:08

I havn't read this whole topic yet so I don't know if this has been talked about or not. I just found out that HNO3 can be formed by oxidizing ammonia. How strong does the oxidizer have to be? also, what glass ware or equipment do I need for this?

Rhadon - 18-1-2003 at 18:41

DeusExMachina, sorry to disappoint you, but it's not that easy. The oxidation of NH3 is an industrial process which is not applicable for use in a small laboratory.

Anyway, I'll briefly describe the reaction.
NH3 is oxidized with the oxygen that is contained in the air at 800 - 900 °C (usually with Pt- or Pt/Rh-catalyst), preferrably in an atmosphere with increased pressure, to NO. The NO is also oxidized by the oxygen which is contained in the air to form NO2, which can then (under increased pressure; still containing unreacted oxygen) be lead through water to yield HNO3 with a concentration of 50-68%.

DeusExMachina - 18-1-2003 at 19:22

sounds complicated to me... have you heard of anyone making it?

Rhadon - 18-1-2003 at 20:01

No. Like I wrote above: It's an industrial process which cannot be carried out in a small lab.

BASF - 18-1-2003 at 21:50


DeusExMachina, sorry to disappoint you, but it's not that easy. The oxidation of NH3 is an industrial process which is not applicable for use in a small laboratory.

No!! It CAN BE MADE in laboratory.

I can tell of at least two experimenting books descrbing relatively simple methods for lab preps.

The most important thing is the catalysator.

This is not that difficult.
Buy 50mg platinum wire(available from chem supplyers quite easily), hammer the wire flat to increase its surface.
Then it is dissolved in hot aqua regia and left for some days til it is completely dissolved.

Then the resulting hexachloro platinum acid is neutralized with ammonia and a precipitate forms, which is directly adsorbed by diatomeous earth.

Finally, the dried catalyst is glowed with the effect, that metallic platinum is left on the surface.

A fire-resistant test pipe is filled with the catalyst and heated to approx. 700°C.

The needed NH3 should be generated relatively dry using some ammonium salt mixed with CaO and the right amount of water.
They warn not to use pure oxygen instead of air, because this can very well result in an explosion.

The two procedures do not cool the reaction gases to increase the yield, as red brown gases can easily be seen though.

Also a Fe2O3- catalyst on the same carrier can be used.
Then the working temp is more than 100°C higher.
Finely divided Fe2O3 can be made by mixing solutions of Fe3(SO4)2 with NaOH and then adding H2O2. :)

Rhadon - 19-1-2003 at 05:51

The procedure still doesn't seem to be applicable for small laboratories to me, but that should depend on the definition of "small".

Thanks for the hint with the catalyst. Do you know of anyone who has tried the whole procedure or even tried it yourself? If so, where the NO2 yields acceptable?

Praetorian - 19-1-2003 at 10:49

Has anybody problems with red fumes NOx during distillation of the anhydrous HNO3 ?(I can't help myself, I think the most effective and simpliest methode is distillation with H2SO4 under suction). No problem. Simply add a few urea into reaction mix. Urea reduce NOx to N2 and carbon dioxide. Excess of nitrous acid is commonly remove after diazotization reaction in a similar way.


a123x - 24-1-2003 at 15:03

Well it has taken me a while bfore I had the time but I finally went out ad got some Ca(OH)2 to make calcium nitrate and some fiberglass cloth to use as a filter so hopefully I'll have results on whether the precipitation method works. Another thing is, don't bother wit oxidation of ammonia for nitric acid producton as it requires the high temperatures and all that crap. Something similar that you might try though is the electric arcmethod of producing NO2. All you'll need is a flyback transforme, an oil burner ignition transformer, NST, or even a microwave oven transformer. You run the arc in a closed container causing NO2 to form and pump this into a cooled container with water vapor being pumped in. HNO3 forms and precipitates. I would have tried this except my flybacktransformer circuit broke and I haven't gotten it working. I think I still might have a working NST and I know I still have a working MOT. That page gives a description of producing NO2 in such a fashion.

Mongo Blongo - 24-1-2003 at 20:10

What do you mean when you say the HNO3 precipitates ?? It just reacts with the water (nothing ppts). It is not an easy method by the way. I have tried it with disappointing results. There is more info in the E&W forum. 10fingers describes his experiences with this method and how to get a higher concentration using water vapor.


a123x - 25-1-2003 at 08:15

I had meant it as the NO2 and water vapor combine and the cooling causes the nitic acid to become liquid from the gaseous state hat it is in. Like precipitation when tlking about the weather is water vapor in the air becomes liquid or solid and falls.

menchaca - 30-3-2003 at 15:15

i dont know what i did when i was child or how i did it but this is what happened to me:

i took a large bottle and i added quantity of ammonia and a large quantity of comercial H2O2(3%) suddenly it´s colour turned red and white fumes appeard in the mouth of the bottle. Could that generate nitric acid? i dont know what to to think because nitric acid isnt easy to obtain oxidying ammnia but this was what happened to me and.........
well that´s all, i hope somebody can help me

PHILOU Zrealone - 30-3-2003 at 23:16

NH3 + O2 may form NOx but usually you need a catalyst!
Maybe there was some unknown catalyst in your ammonia solution or in your H2O2 solution!

osamafon - 31-3-2003 at 05:52

can I make 95% hno3 +NH4HSO4 with NH4NO3 and H2SO4 ? if it is posible ,is it good for RDX or PETN?

johnybmb - 1-4-2003 at 03:40

You could try to centrifuge the mix and decant most of the nitric acid , as the calcium sulfate would not, for a short time , float on the dense nitric acid.

PHILOU Zrealone - 1-4-2003 at 17:12

"can I make 95% hno3 +NH4HSO4 with NH4NO3 and H2SO4 ? if it is posible ,is it good for RDX or PETN?"

Yes if your H2SO4 is 95%!
RDX needs H2SO4 free HNO3 or a minimum since H2SO4 destroys fast RDX!

PETN procedures don't use H2SO4 but rather use HNO3 conc in excess, HNO3/P2O5 or HNO3/Ac2O; so I guess there should be a problem with H2SO4!

Better separate HNO3 from H2SO4 and NH4HSO4...Use less than stoechiometric amount H2SO4 to keep a slight excess of NH4NO3; cool down the mix and filtrate the NH4HSO4 cristalls!The excess NH4NO3 will help precipitation/ristallisation of NH4HSO4!


Cappy - 4-4-2003 at 08:51

Apparently HNO3 can be dehydrated to greater than 68% concentration by hydrating Mg(NO3)2. "For ease in handling the material it is highly desireable that the magnesium nitrate be in liquid form." This means dissolving it in water, since it melts at too high a temperature. Since water is added too the HNO3, dehydration is less effective.

Would adding Mg to HNO3 work? Sure it would use up some of the acid, but Mg is sometimes easier to find than Mg(NO3)2.

PHILOU Zrealone - 7-4-2003 at 03:51

It is a good idea, but most of the people scream for Mg and would consider this as a crime

AngelEyes - 7-4-2003 at 07:17

Seems to me the best way to dehydrate HNO3, if you're going to do it by adding a dehydrated salt, would be to use CuSO4. You know when it's anhydrous 'cos it goes white, you know when it's absorbed the water 'cos it goes blue, it's cheap and (should be) readily available and it won't react with the Nitric acid either.

I think that's all correct - unless anyone knows different? I haven't actually tried it myself though...


dynamite - 10-4-2003 at 01:56

is the reaction with 2 nh4no3 +h2so4 so hno3+nh4hso4+nh4no3. ??? if is how can I do it so that 2hno3 forms , with a cons. acids(96% h2so4)

PHILOU Zrealone - 10-4-2003 at 06:13

2NH4NO3 +H2SO4 --> HNO3+NH4HSO4+NH4NO3
Is what you wrote!

Then simplifiying NH4NO3 in the two parts of the equation, you end up with:

NH4NO3 +H2SO4 --> HNO3+NH4HSO4

So simply add equimolar amounts of H2SO4 and NH4NO3 to get molar amount of HNO3!


Cappy - 14-4-2003 at 16:58

Could you produce nitrogen dioxide by heating air with a lightbulb? It would then be dissolved into water to form nitric acid. How long do you think the tungsten would last in air, and would it contaminate/catalyse the reaction at all?

The air would be pumped in somehow (which would explain why the water is not travelling up the glass tubes).

[Edited on 4/15/2003 by Cappy]

trinitrotoluene - 14-4-2003 at 19:28

I heard there was some way to produce HNO3 from ammonia. 2 NH3+ O2= 2HNO3 or do you need some sort of cataylist?

Madog - 15-4-2003 at 03:45

cappy, i dont think that would work, at least not well, electrolisis of air is how to do it. and add some pure oxygen to try to get it to a good mix of N2 and O2.

edit: oh yeah, also, the tungsten is not going to last long in air at all. it will oxidise as soon as you put current through it.

i might electrlisis some air just for the hell of it, i got a transformer that can do the job.

[Edited on 4/15/03 by Madog]

Cappy - 15-4-2003 at 21:40

Is there a minimum voltage or current needed to electrolyse air? Also, should the electrodes be very close, or as far apart as possible while still allowing a visible spark? Is DC required or will AC work? I believe AC would work if just heating was required.

[Edited on 4/16/2003 by Cappy]

madscientist - 15-4-2003 at 23:08

The purpose of the electric arc isn't really to electrolyze air - it's to create temperatures so extreme that the triple bond in nitrogen molecules breaks, allowing nitric oxide to form (which is then oxidized to nitrogen dioxide). That in mind, I would go with the longest arc possible, highest voltage, and alternating current.

a123x - 26-4-2003 at 17:58

I just had a thought. Since nitric acid forms an azeotrope with water at 68% and this azeotrope boils at 20C higher than water boils at then nitric acid of concentrations below 68% ca be concentrated to 68% by simply boiling off water right? This makes me think the Ca(NO3)2 + H2SO4 method could actually be quite good if the nitrate was dissolved in water then the H2SO4 added. All the water would inrease the volume quite a bit making decantation quite easy. Then the acid could at least be concentrated to 68% which is good for most nitrations and should be better than H2SO4 + KNO3 for those nitrations.

a123x - 26-4-2003 at 19:07

I had an even better idea. A fiberglass cloth filter is plenty good enough to stand up to acids and if the strands are bunched together by hand then it is tight enough to prevent even quite fine particles to pass through it. The problems becomes that the CaSO4 takes up quite a bit of volume so that there may not even be any free flowing liquid. I have thought of a way to solve this problem. Take a strip of fiberglass filter, fold it over, take some fiberglass epoxy and put it into the sides of the folded strip so that whe dried it forms a pouch type thing with only one end open. This allows you to take the CaSO4 + nitric acid mixture and pour it into this pouch, then use some type of plastic roller made from something like HDPE which has pretty decent resistance to nitric and roll the pouch with it from the open end down to effectively squeeze out the acid.

Madog - 29-4-2003 at 17:01

cool, you could just smash open the wall and pull the pink stuff out and use this in vacume filtration.

[Edited on 4/30/03 by Madog]

Haggis - 30-4-2003 at 07:45

Attics are the best source for the stuff. There is a whole lot just hanging on the walls and floor. Some houses use spray type insulation and I'm not sure that this would work because it feels completely different from the pink stuff. They consistency is not uniform with the thickness or concentration of the material. The pink stuff would be the best choice, if it does indeed work. If you cannot get it from your own attic, go to a construction site where they are building a house. I work construction and there is always 'pink stuff' lying around. The best times are when they first put the wall frames up. There is usually a layer between the top of the basement and the bottom of the floor. The second best time is when it is starting to look like a house. The walls have plywood and the roof is complete, before they put on siding however. It is layed between exterior wall studs. They usually don't insulate between interior rooms however. This is some itchy stuff, so wear long sleeves and some type of glove. Just ask the guys for some insulation scraps. I find the brand with the "Pink Panther" on the roll to be the highest quality.

leopard - 8-5-2003 at 13:02

I make my Nitric Acid using H2SO4 + KNO3 (equal proportions by weight) but my yield seems very low. In Rudolph Wagner's "Handbook of Chemical Technology 1872" he states that typical values for materials are 30kg KNO3 and 29kg H2SO4 then goes on to state that typical yield is 125kg Nitric Acid for 100kg KNO3 !!! Molecular weight of KNO3 is (101) while for HNO3 is (63). I'm no chemist but the math doesn't seem to add up. If the K (39.1) is taken away and replaced with H (1.0) how does the final weight of HNO3 exceed the starting weight of KNO3 ? Am I missing something obvious? By the way I also have good success making HNO3 using 2 neon transformers each creating a 1.5cm arc inside seperate jars. A pump forces air into the first jar then into the second jar then bubled into water.:)

Marvin - 9-5-2003 at 19:05

The reasoning is right, but you didnt read the section all the way through. The process described condenses the nitric acid into water, producing an acid of concentration 36 degrees Baum.

This is around 50% nitric acid. The additonal discrepancy from the numbers comes from the use of sodium nitrate, not potassium nitrate to make the acid.

Watch out for NO2, its very toxic, avoid breathing it in at all costs, at low concentrations it will cause lung damage while you dont even notice it. If I sound overly repetative on this point, its becuase NO2 bit me back.

Reedited after working some things out.

[Edited on 10-5-2003 by Marvin]


Theoretic - 24-6-2003 at 07:18

I heard that a mixture of N2, O2 and Cl2 forms NOCl (nitrosyl chloride) on heating to 400C. You could then use sunlight to decompose NOCl to NO and Cl2 (chlorine functions as a catalyst for fixing nitrogen).On air NO is oxidized to NO2, which could be liquefied and separated from chlorine, that could then be reused.
Nitrogen and oxygen from air would do just fine.:D

a123x - 25-6-2003 at 08:35

That would be way more trouble to do than it is worth. For one thing you're working with Cl2, then you have to liquify NO2 to get it out of the system. It seems it would be significantly easier just to run a neon sign transformer in a container with air being slowly pumped in one side.

Theoretic - 26-6-2003 at 07:08

Liquefying NO2 isn't hard, it's boiling point is 21.1C (this has been taken from a Material Safety Data Sheet for N2O4 - liquefied NO2).

VoD - 16-7-2003 at 05:33

If I were to use the conc. H2SO4/KNO3 method with an oil bath and a pyrex distillation setup (a condenser), then distilled at about 65C so no H2O would evaporate (I think 65 is the evap point of HNO3), would I have 100% HNO3 (or close to 100%)? Would the final product be sufficient for RDX manufacture?

a123x - 16-7-2003 at 10:15

You should use a bit of excess H2SO4 since it has high hydroscopicity and will hold any water from going through the condensor but really you should be just using 98% H2SO4 from the start so there is barely any water anyway. Keeping the temp at 65C is good to do because it will help prevent the HNO3 from decomposing although I think you could take the oil bath to 70C. If you have to ask these questions about HNO3 distillation then you're not ready to make RDX.

VoD - 16-7-2003 at 10:25

Oh, I am ready to make RDX, I have been doing this stuff for many years, it is just that I was reading up and some people are saying that a simple distillation won't yield anywhere near 100%. Since >95% is necessary for RDX, I dont want to waste my time and money on the glassware if it will be <95%, as nothing else I am planning to do will require >70% (which I can just buy)...
Do you think I should spring for the glassware (condenser, teflon stoppers, ect.) or not bother? RDX is my main goal with this and if its no dice on >95% HNO3 then I won't bother spending the money.

a123x - 16-7-2003 at 12:39

I think what they meant is that normal distillation of HNO3 can't yield a concentration higher than about 70% because it forms an azeotrope with water at 68% HNO3 so that you can't get higher than that unless H2SO4 is used to keep the water from evaporating with the HNO3. A proper condensor and teflon stopper isn't necessarily needed. I've recently had success making nitric acid of concentrations of around 95% with a very improvised still. It's a large glass jar/beaker with H2SO4/NH4NO3 mix in it, in the center of that is a glass or other acid resistant stand(I used a small glass jar), on the stand is a small beaker/glass jar. The opening of the large jar is covered with a HDPE sheet sealed on with a rubberband. Ice is placed on the sheet with it loose enough to dip down in a bowl shape. For mine I actually use part of a light fixture that is like a glass bowl with a conical bottom. Anyway, this is put in a hot water/oil bath kept at 60-70C. I've used my set up three times giving me like 95% HNO3 each time and about 7.5-10g per 2 hours. Look on brain fever's site for some pictures of a still like this. Also you could get some teflon tubing on ebay to make a condensor, I just recently bought 250ft of 1/8" inside diameter to make a still with.

VoD - 17-7-2003 at 05:09

Thanks, I may have to try that... or just buy the apparatus off of ebay....
Have you had success using your HNO3 to make RDX?

[Edited on 17-7-2003 by VoD]

a123x - 17-7-2003 at 12:59

Haven't tried RDX yet. Actually I'm just planning to do my first nitroaromatic in the next few days. So long as you use 98% sulfuric acid with the KNO3 to distill the nitric from it should be 95%+ nitric and pure enough for RDX.

VoD - 17-7-2003 at 14:57

I have 98% H2SO4, however it is drain cleaner so it is dyed.... that shouldn't matter though I suppose. I plan on using KNO3 instead of AN, so I guess I will do the calculations and give 'er. I don't like how ghetto brainfever's setup is though, so I may try something a little different though...
How do you plan on using the teflon tubing for a still?

a123x - 17-7-2003 at 17:10

I intend to take a flask and stopper it with a teflon stopper(my uncle bought a lot of plastic including a bunch of teflon from a distributer that was going out of business, he also has the tools to machine the stopper). Then I'll have one hole in the stopper with the tube sticking directly in it, from there I will wrap the teflon tube in a single layer several times around an ice-filled bottle, after that the tube will just drip into another flask. Considering that the tube is only 1/8" internal diameter I might have like 5 holes in the stopper for five tubes and just have them wrapped side by side around the same bottle of ice so I can condense more at a time. Heat will still be with hot water or oil kept at around 70C. I made a quick illustration of it here:

BromicAcid - 18-7-2003 at 18:17

Calcium oxalate is insoluble as all hell, maybe Ca(NO3)2 + H2C2O4 ---> 2HNO3 + Ca2C2O4 but I dont know how the oxalate would fair in such an increasingly acidic enviorment. Another worry would be Oxalic acid being precipiated out before the reaction is anywhere near complete due to it being a weak acid and the equilibrium being shifted due to increasing pH of the solution, anyways, it's just a thought, maybe more manageable then the CaSO4 precipiate, maybe a constat system at high temperature could be done with slow addition of the nitrate to a oxalate solution with effient boiling off of the nitric acid as it is formed keeping equilibrium favored.

Mumbles - 18-7-2003 at 22:53

This is just a thought I've been having. Whenever people do the precipitation method they add all the Calcium nitrate at once and can get none or little Nitric acid. What if you added it in portions. Add an amount, then filter. Add another amount and filter again. This would increase the amount of nitric acid everytime. I'd imagine that a smaller amount of precipitate could be pressed harder to remove the remaining acid as well.

This has been running through my mind for a couple days. I've been trying to figure out what wouldn't work. I can't find a reason for it not to. I know its far from perfect, but it may be a way to get a signifigantly larger amount of very high purity Nitric acid in a relitively small amount of time.

Theoretic - 23-7-2003 at 07:58

How about this:
1) 2NH4NO3+Ca(OH)=>2NH3+2H2O+ +Ca(NO3)2

BromicAcid - 23-7-2003 at 17:13

I don't know about that decomposition of calcium nitrate, usually nitrates decompose along the lines of

Ca(NO3)2 ----> CaO + O2 + 1/2N2

although I know some nitrates favor decomposition to nitrogen dioxide such as lead nitrate.

Pb(NO3)2 -----> PbO + NO2

This reaction is very reliable and easy with the exception that the lead monoxide attacks glass so take some aluminum foil and use it to line your flask, any oxygen or nitrogen impurities should not bother your dissolution into the water so no worries there, this reaction gives a good yield and is quite controllable. Supposedly it decomposes around 470 C. That actually seems a little high, maybe it was basic lead nitrate... 3PbO*N2O*H2O nahhh, the formula makes it seem like it would let up that water molecule and the molecule of nitrous oxide, regardless, I've got a nice jar of it in my lab from the 1800s that reads [Plumbous Nitrate For Making Nitrogen Tetroxide] goes right along with my jar of iron sulphide, notice the p in sulfide, good ol' vintage chemicals.

Oh yeah, and this:
2NO2 + H2O ---> HNO3 + HNO2
but nitrous acid decomposes easily expecially if you elevate the temperature a bit
3HNO2 ----> HNO3 + 2NO + H2O
from here you could react the nitric oxide that comes off with more oxygen and put it back into the system. One more thing, there is an equilibrium
2NO2 <----> 2NO + O2
as the temperature rises it gets more and more toward the NO side and if oxygen gets out of the system they will not recombine and you'll end up bubbling mostly NO though your water unless you have an efficent condenser. Like I said, the decomposition reaction above seems a little high but if it is running that hot I'm glad I've got my condenser running.

[Edited on 24-7-2003 by BromicAcid]

blip - 23-7-2003 at 22:06


Ca(NO3)2 ----> CaO + O2 + 1/2N2

Small error: ;)
Ca(NO<sub>3</sub>;)<sub>2</sub> <sup><u>&nbsp;<font face="symbol">D</font>&nbsp;</u></sup>> CaO + <sup>5</sup>/<sub>2</sub>O<sub>2</sub> + N<sub>2</sub>

Thanks for the great bit about Pb(NO<sub>3</sub>;)<sub>2</sub> <sup><u>&nbsp;<font face="symbol">D</font>&nbsp;</u></sup>> PbO + 2NO<sub>2</sub> + <sup>1</sup>/<sub>2</sub>O<sub>2</sub> :)
I wanna liquefy N<sub>2</sub>O<sub>4</sub>. :D

rikkitikkitavi - 23-7-2003 at 23:05

sorry, big error:
where did you get the idea that nitrates decompose into N2?

Almost all metall nitrates decomposes into some nitrogen oxide. The heavier the metall nitrate, the more favour of NO2 :) and generally lower decompostion temperature too (depends of course on the stability of the nitrate ) Lead and copper almost entirely decomposes into NO2 and metall oxide.

Since Cu (NO3)2 formes a hydrate with low melting point , you will actually distill of HNO3,H2O, NO2 from it when heated !
Not nice when you dont know about it ....
cough, coough...

Ca(NO3)2 decomposes into NO2 ,NO, N2 and O2

rougly Ca(NO3)2 => CaO + 2 NO2 + 1/2 O2 starts at around 500 C, melting point for CaNO3 is 560 C, where decompostition is vivid.

But because of the high temperature much of the NO2 decomposes into NO + O2, but later when temperature is reduced , NO2 forms again.

All taken from Gmelin and Ullmann. I havent tried in real life, but when Ca(NO3)2 is heated strongly there is a faint brown gas emitted, sofar I have come. It is time to build a retort

[Edited on 24-7-2003 by rikkitikkitavi]:D:D

[Edited on 24-7-2003 by rikkitikkitavi]

Theoretic - 24-7-2003 at 09:46

So it WOULD work after all...:cool:

BromicAcid - 24-7-2003 at 18:37

Sorry about that, never decomposed calcium nitrate by it's lonesome, I'm used to working with nitrates in pyro and therefore with a good fuel you get pretty much all oxygen probably because any NO2 formed is going to be an active oxidizing agent. I was sure about the lead nitrate though. Anyways, in the future decomposition of nitrates will be duely noted as being not as straight foreward as I would like to believe.

rikkitikkitavi - 25-7-2003 at 00:17

according to litterature it will work, but wether it is possible IRL we have to do experiments. Perhaps yield isnt so good.


Theoretic - 31-7-2003 at 08:35

OOPS... the problem is where we least expect it (more so if you need anhydrous HNO3).
4NO2+2H2O+O2=>4HNO3, eh...
Well if it's 70C, 50atm and 4 hours (for anhydrous HNO3)!

HNO3 Percentage vs Density

chemoleo - 3-11-2003 at 07:42

Has anyone got a detailed chart for the percentage of HNO3 vs it's density? Often they quote densities but not the percentage... any sources anyone?

% in what? H2O?

PrimoPyro - 3-11-2003 at 10:37

What are you looking for? The density of x solution of HNO3 in water? For water it's easy. Calculate it. You know the ratios of the components of the solution,the weights, volumes, and densities,of both.

Calculation time. :)

well well... might not be that easy!

chemoleo - 3-11-2003 at 12:22

well, of course it's H2O. If it was another solvent (like what exactly? D2O , DCM, H2SO4?) I would have mentioned it.

I thought that one might calculate it, however, it might be possible (though I dont know) that the density of a mixture of solvents increases/decreases non-linearly with increasing one over the other.
What I mean by this is that two different solvents of different densities don't necessarily behave according to

density_final =
(density1 x volume_solvent1/volume_total) + (density2 x volume_solvent2/volume_total))

when mixed.

For instance, the density of solvent1 is 0.5, and that of solvent2 is 1.5. I take 50 units (teaspoons , gallons, ml ;) ) of each. Thus you would immediately guess the final density is 1. Using that equation works of course. I.e. density_final = (0.5x50/100)+(1.5x50/100) = 1.

However, I am sure there are cases where the solvent spaces (i.e. spaces between the two types of molecules) enable the final density of the two solvents to go up (or down?), that is the final volume decreases (I guess this depends on molecular size, shape, van der Waals & electrostatic forces, and even chemical reactions such as H2SO4+H2O --> [H3O+] + [HSO4-]) so it would be a solvent-specific system, hence not predictable) ! In the above case, 50 ml of solvent 1 plus 50 ml of solvent 2 would be LESS than 100 ml when mixed.

So, how do you know this is not the case, primopyro? See, that's the reason I was asking :)

So again, anyone got a % vs density chart of HNO3 in H2O? :P

Edit: In fact I just remembered that I did this calculation once with HNO3 and a few known densities/%'s. Surprise surprise, the theoretical value was DIFFERENT from the real one, kind of making a chart necessary!

PS Primo I am waiting for an answer! ;)

[Edited on 2-12-2003 by chemoleo]

vulture - 2-12-2003 at 01:34

A good example is water + ethanol;
mix 50ml of water and 50ml of ethanol and you end up with 96ml of mixture.

Density Chart of HNO3

chemoleo - 5-12-2003 at 16:35

Finally found it, in the roguesci archives. For anyone curious, yes, it is HNO3 in WATER, not D2O etc :D :D

w[%] d[g/ml]
0 1.0000
1 1.0036
2 1.0091
4 1.0201
6 1.0312
8 1.0427
10 1.0543
12 1.0661
14 1.0781
16 1.0903
18 1.1026
20 1.1150
22 1.1276
24 1.1404
26 1.1534
28 1.1666
30 1.1800
32 1.1934
34 1.2071
36 1.2205
38 1.2335
40 1.2463
42 1.2591
44 1.2719
46 1.2847
48 1.2975
50 1.3100
52 1.3219
54 1.3336
56 1.3449
58 1.3560
60 1.3667
62 1.3769
64 1.3866
66 1.3959
68 1.4048
70 1.4134
72 1.4218
74 1.4298
76 1.4375
78 1.4450
80 1.4521
82 1.4589
84 1.4655
85 1.4689
86 1.4716
87 1.4745
88 1.4773
89 1.4796
90 1.4826
91 1.4842
92 1.4873
93 1.4886
94 1.4912
95 1.4932
96 1.4952
96.5 1.4972
97 1.4988
97.5 1.5005
98 1.5008
98.5 1.5044
99 1.5066
99.5 1.5091
100 1.5129

See also the Calculation Concentration thread for more ....

[Edited on 6-12-2003 by chemoleo]

There is a better way!

pyroscikim - 21-12-2003 at 19:41

If you could get methylene chloride lucky you! cos you could make >98% nitric acid easily.

make a solution of 45mL 98%sulphuric acid and 55mL 70% nitric acid, let cool to room temperature and add 100mL methylene chloride. Stir for 2minutes, let the liquid settle and extract the methylene layer into a sealed container. Pour into the acid another 100mL of methylene chloride and repeat for up to 8 times. The nitric acid will form a complex with methylene chloride but water and sulphuric acid will not dissolve at all.

when all 8 extractions are complete, distill the methylene chloride solution to yield methylene chloride at the condenser and pure HNO3 is left behind. this is easy since methylene chloride boils at 41degrees Celcius. yield is about 80% of the nitric acid (not taking into acount the water) began with.

the sulphuric acid can be recycled if wanted by heat-concentrating. the methylene chloride is used for another batch.

Where i live (australia) i could find a solution of 87% methylene chloride in methanol, used as a remover for dry paint, and i believe you could extract pure methylene chloride by adding equal amount of water to the methylene/methanol solution, and extracting the insoluble methylene chloride.

BromicAcid - 8-9-2004 at 17:42

I don't know about that decomposition of calcium nitrate, usually nitrates decompose along the lines of

Ca(NO3)2 ----> CaO + O2 + 1/2N2

I know rikkitikkitavi corrected me further up thread but I was reading "The Chemistry of the Elements" and under the entry for calcium nitrate it lists the normal decomposition of it. If heated slowly first:

Ca(NO3)2 ----> Ca(NO2)2 + O2

Followed by if heating is continued:

Ca(NO2)2 ----> CaO + NO + NO2

However if heated quickly it would probably pass right though that nitrite stage and go to the total decomposition giving oxygen and nitric oxides and just as rikkitikkitavi said they probably dissociate readily at this temp but upon cooling you would be left with a large percentage of nitric oxide.

(Just kept bugging me after I read it, took me 30 minutes to find this thread.)

Edit: Hummm... 75th post to this thread and therefore according to my viewing settings my own page, not good, now I'm not in context with the rest of the thread and my point is partially lost. Basically this was just an example of me correcting myself and agreeing with another person.

[Edited on 9/9/2004 by BromicAcid]

Theoretic - 20-10-2004 at 01:00

Sorry, about me scaring you with high temperatures, pressures and long times for NO2 + O2 dissolution, we don't necessarily want 100% acid. :)
A way to separate your HNO3 from the precipitate is to distill it with the plastic-film-over-vessel-and-beaker-underneath method described for production of HNO3 from NaNO3, KNO3, NH4NO3 and the like (distillation below boiling point - DBBP).
As I found out from someone's link, NH4NO3 can be decomposed catalytically by Cl- to give dilute HNO3 vapour.

Quince - 15-9-2005 at 10:47

Can the methylene chloride method be used to extract directly from the H2SO4/nitrate mix usually used for distillation of HNO3? Or will the bisulfates interfere?

Eclectic - 15-9-2005 at 11:06

Yes, there were some aromatic nitrations using CH2Cl2, NaNO3, and an acid on Rhodium's site. I don't have the paper handy, but I think the PDF is on FTP2 in the Rhodiums PDF folder.

What is the problem with distilling 100% HNO3 from an equivolume mix of 98% H2SO4 and 70% HNO3? That's how I did it back in my college days for a prep of tetranitromethane from acetic anhydride and 100% HNO3. (back in olden times before everything caused cancer and chemicals were deadly national hazards)

[Edited on 15-9-2005 by Eclectic]

Quince - 15-9-2005 at 17:27

The problem is I don't have any dilute nitric acid; all my HNO3 this far has come from H2SO4/nitrate vacuum distillation. I'm trying to figure out if I can skip the distillation, as the aspirator uses up an immense amount of water.

I don't intend to nitrate with the mix. The dichloromethane is supposed to evaporate faster at a lower temperature than the HNO3. I could just put a vigreux column with an open top, and the flask in 50*C water bath.

Eclectic - 16-9-2005 at 05:03

68% is the azeotrope you get when you distilling nitric acid at atmospheric pressure from a mix of dilute H2SO4 and a nitrate. Just use your vigreux column without vacuum.


[Edited on 16-9-2005 by Eclectic]

quicksilver - 16-9-2005 at 06:31

Originally posted by Quince
The problem is I don't have any dilute nitric acid; all my HNO3 this far has come from H2SO4/nitrate vacuum distillation. I'm trying to figure out if I can skip the distillation, as the aspirator uses up an immense amount of water.

I don't intend to nitrate with the mix. The dichloromethane is supposed to evaporate faster at a lower temperature than the HNO3. I could just put a vigreux column with an open top, and the flask in 50*C water bath.

You -=CAN=- skip the vacume. I have tested this for some time with some fairly good glassware and I get the same very, very light, lemon yellow acid with a vacume as without. I went up to as high as vacume as I could pull and it stayed the same as no vacume (faintest yellow) and all the labs utilizing that HNO3 were just fine.
On another note, I used to distill dichloromethane from a paint removed and get quite a bit of the pure product. When I use dichloromethane I can receive the HNO3 from it via distilltion at 40 C (no higher) thus I use a "hot pad" (for beds, bad back, etc); it keeps the temp at a perfect 104F and my distilltion is very clean and fast. I have stopped using a vacume entily as I had access to a really good pump and I just couldn't get 100% clear HNO3 without I just use the Urea and live with it. It really does work fine.

12AX7 - 27-9-2005 at 09:49

Found some impure nitrate in the basement (Osmocote or whatever, bulb feed, contains ammonium nitrate). I put some salt (I filtered it from the CaPO4 and organic casing parts and recrystallized it, if you can call it crystallized) in the bottom of a jar with H2SO4 with a smaller jar supported inside, with LDPE plastic covering it with ice. It's frothy and the vapor is brown at the moment... :o I've collected an ounce or two of distillate so far but it doesn't react with copper metal like HNO3 is supposed to...

Edit: Of the distillate I've collected so far, it does not react with copper but does produce a lot of white precipitate with Cu2O. :(

The apparatus at the moment has quite a bit of orange color to it, but last I checked not a drop was condensed. What the heck?


[Edited on 9-27-2005 by 12AX7]

IPN - 27-9-2005 at 10:50

Redistill from H2SO4. It's probably just too weak to dissolve copper.

Condensing Nitric Acid

Cyroxos - 30-12-2005 at 15:22

Lotek and I were trying to create some Nitric acid. We used about 150ml of H2SO4 and about 250 grams of KNO3 in a round bottom flask and preheated the acid to just over 100 degrees C. We were expecting to condense the red Nitric acid; yet, the colorful gas only stayed in the flask and did not move through the condenser coils. As shown in this picture.

The liquid in the flask was boiling, and there were not blockages in the flow to the condenser. The stoppers were sealed. We waited, yet, the (what I assume to be Nitric acid) liquid condensed on the sides of the flask and the glass rods, yet never condensed in the condenser. Are we just not waiting long enough? Do we need to apply more heat to it? What exactly are we doing wrong? What about the small diameter of the glass rods? could that be a problem

- Cyro

lacrima97 - 30-12-2005 at 15:36

I have had this same problem using a retort. The only advise I was given was to crank up the heat, and then bubble O2 through the NO2 contaminated HNO3, to make it clear..

chromium - 30-12-2005 at 15:41

Not enough heating or maybe you not waited enough time. Brown gas does not condense, after some time most of gas will be colorless and invisible. This can be condensed to get nitric acid. Product will be red (or brown) as brown NO2 dissolves in nitric acid.

[Edited on 30-12-2005 by chromium]

Magpie - 30-12-2005 at 15:45

That brown gas is NO2 and its presence is normal when heating nitric acid.

Water/HNO3 should come off first if your mix is rich in water. When the pot concentration reaches 68% HNO3 (azeotrope) this should come off steadily at 120C until you are done.

You would have to be initially lean in water to get to just fuming nitric acid. Fuming nitric may require vacuum conditions to prevent decomposition. I have not made fuming HNO3 so can't say from experience. I only made the 68% acid.

I think you are just not waiting long enough. I used a straight tube condenser allowing a clean cut when the temperature reaches 120C. It seems like your helical condensor will have some holdup but it should do the job.

Search "Making own nitric acid" and you will pull up some experience that should be useful.

[Edited on 31-12-2005 by Magpie]

[Edited on 31-12-2005 by Magpie]

[Edited on 31-12-2005 by Magpie]

The_Davster - 30-12-2005 at 16:05

Rubber stoppers will be eaten by nitric fumes. Wrapping your stoppers in saran wrap will stop this.

Your setup is not well suited for a distillation of something as sensitive as nitric. The small tube connecting your boiling flask to the condenser is rather thin, so it will air cool and mainly act as a reflux condenser, making you need higher heat to get the nitric fumes into the condenser, and higher heat will also decompose the nitric. Try getting a larger diameter glass tube for connecting the two, and or insulate the tube to prevent condensation here. Glass wool or cork will work for this. Insulating the top of the flask the same way would also be beneficial. For nitric, a nice, slow distillation gives the best product.

Magpie - 30-12-2005 at 16:15

I agree with rogue here that this synthesis is nothing to take lightly. My rubber thermometer adapter did suffer some degradation due to the nitric acid fumes. A keck clamp at a leaking ground glass joint was destroyed. Both NO2 and HNO3 are not something you want to breathe. I was very glad I had good ventilation (a hood).

If you are just after fuming nitric acid you will need vacuum conditions (I believe) which presents a whole other level of required equipment sophistication and experience. Just imagine a faulty RBF imploding full of nitric/sulfuric acid and you are standing in front of it without adequate protection. It would not be a pretty sight.

Aluminum foil will also act as insulation.

kABOOM! - 30-12-2005 at 17:09

...or my way of getting fairly decent grade Nitric Acid...never mind about the fancy condensing coil method...just take two pyrex flasks with stoppers and a long 3/8" pyrex tubes and have 1 of the flasks mixed with Nitrate/ H2SO4 and the other beaker empty in an ice bath slightly tilted downwards. I made 800ml of fairly good grade Nitric acid this way. It was nearly colorless with only a very slight yellow tint to it. You get Nitric acid very quickly this way because the air is able to cool the gases much faster using a larger diameter glass tube. I'm not saying that the condensing coil method doesn't couse not. It just takes a much longer time--I don't find it nearly as effective.

The_Davster - 30-12-2005 at 17:30

Also, a condenser like that works best in a vertical position, so that the nitric that forms drips out of the condenser and does not stay in the loops.

What is the other tube leading into your condenser for?

Magpie, fuming nitric does not actually require a vacuum setup, I have made small ammounts(<10mL, enough for my purposes) using only a 600mL beaker, a shot glass, and a plastic baggie. The method is on Brainfevers site.

neutrino - 30-12-2005 at 21:35

I have made fuming nitric acid by atmospheric distillation without problem. I used the standard mix of AN and SA (with a little extra SA), a 500mL RBF, 400W fiberglass mantle, and 300mm liebig, all with 24/40 ground glass joints. The acid came over with minimal decomposition, certainly far less than what is in that photo. It is yellow, I can take a picture if anyone wants to see it.

Someone else here did this a few months ago and concluded that decomposition is mainly from sunlight and that the whole vacuum thing is just a bunch of bogus science. I don't remember who it was, but I have to agree.

kABOOM! - 31-12-2005 at 00:44

I often use AN with H2SO4 and I get fuming Nitric acid. I've made all kinds of fun energetic materials this way. Right now I have 2.5 L of fuming Nitric acid on the brew. You do need a safe environment to do this in.. not inside a garage with flammable solvents and liquified petrolium tanks. BAD BAD BAD idea. Do your distillation inside a fume hood or outside!!
NOx- Nitric fumes are DEADLY DEADLY... your lungs will pack up in seconds and you'll die a horiffic death.

I nearly died this way many many years ago... never ever do a distillation of Nitric acid in an enclosed environment!

[Edited on 31-12-2005 by kABOOM!]

neutrino - 31-12-2005 at 08:34

>I nearly died this way many many years ago...

Really? Do you have any chronic health problems as a result?

kABOOM! - 31-12-2005 at 12:49

I was lucky to escape serious perminent injury. I suffered some damage to my lungs as a result of Nitric fumes--it took the better part of 2 months to get a clean bill of health. It was much like having a bad lung infection and every breath hurt. I'm 100% better now. Very little scaring--nothing that will cause serious problems in the future. I'm REALLY lucky...the doc told me that many people suffer long lasting health effects due to errosion of the lung tissues. So far so good...I learned my lesson.

Fulmen - 1-1-2006 at 07:17

I have to agree with Rouge that the setup is less-than-perfect, using larger (and shorter) glass tubes and a vertical condenser might improve the output.
I have distilled strong (95% +) nitric acid with wery little discoloration under atmospheric tressure, so vacuum is not nessecery. I used equal parts (by volume) of 65% nitric acid and 98% sulphuric acid and got pretty much the expected yield (half the volume of the 65% NA). The setup was a 1l EM-beaker with taper joint, a vigareux-coloumn (probably not required), a three-way adapter and a fairly simple double-walled condenser.

Cyroxos - 1-1-2006 at 16:24

Yea, we learned that that setup had some serious problems. For one, the rubber stoppers dissolved and filled our condenser with a thick stopper gel... We got that cleaned up though. So, I suppose in time we will do it again differently.

How important in temperature control in this synthesis? We heated the H2SO4 up to 100 C then added the KNO3 and let it go from there without additional heating. We have no way of measuring the temp inside the vessel without taking the stopper off, so we are pretty much guessing and hoping that the temp stays in the 100-120 range.

[Edited on 2-1-2006 by Cyroxos]

The_Davster - 1-1-2006 at 17:14

Having never distilled the azeotropic acid, only the fuming, I used a waterbath at around 80C for a nice slow distillation.

That also explains why you did not get any distillate last time, you did not heat it, only relied on the initial H2SO4 temperature. Seeing as you don't seem to mind getting the red acid, just use a boiling water bath heated by whatever your heatsource is and by also heeding the suggestions above you should definatly get product.

[Edited on 2-1-2006 by rogue chemist]

Dream of the iris - 1-1-2006 at 17:30

How do you calculate the density of your final product? How do you know if you have 70% or 100%?
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