I haven’t tried teflon. For distilling the liquid it’s worth the investement for a distillation flask or better glass retort. For stoppers, I have rubber stoppers I’ve wrapped in Al foil and used that to distill fuming HNO3. The foil gets only slightly to barely attacked. Better would have been a glass stopper.

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Concentrated nitric acid used in laboratory work is 68% nitric acid by mass in aqueous solution. What should be the molarity of such a sample of the acid if
the density of the solution is 1.504 g mL-‘?​

Concentrated nitric acid used in laboratory work is 68% nitric acid by mass in an aqueous solution. This means that 68 g of nitric acid is dissolved in 100 g of the solution.

Then, number of moles of HNO3 = 68 \/ 63 mol

Also density = 1.504g\/mL-1 (given)

Therefore from the formula density = mass \/ volume,we get

Volume of solution = 1000\/1.504 = 66.49 mL

Therefore molarity of nitric acid = (1.08\/66.49) x 1000 = 16.24 M «>,<"id":83860167,"content":"

percentage>[\/tex] (%w.w) = 68 % (w\/w)

solution>[\/tex] = 1.504 g.ml[tex]^<-1>[\/tex]

1.504>< 63 >[\/tex] = 16.23 M «>]» data-testid=»answer_box_list»>

Explanation:

Concentrated nitric acid used in laboratory work is 68% nitric acid by mass in an aqueous solution. This means that 68 g of nitric acid is dissolved in 100 g of the solution.

Then, number of moles of HNO3 = 68 / 63 mol

Also density = 1.504g/mL-1 (given)

Therefore from the formula density = mass / volume,we get

Volume of solution = 1000/1.504 = 66.49 mL

Therefore molarity of nitric acid = (1.08/66.49) x 1000 = 16.24 M

What metal should be used to make a vessel to store concentrated nitric acid

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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.

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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

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?

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( 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]

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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
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Performance

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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. .

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.

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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]

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[Edited on 27-4-2010 by avi66]

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franklyn, its a fact that nitric acid decomposes above 86c, so how this reactionkhso4 + 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.

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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.

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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.

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…

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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]

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

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.

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

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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.

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

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What metal should be used to make a vessel to store concentrated nitric acid

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I was thinking about copper nitrate 2.5 hydrate as a precursor to fuming nitric acid via the normal acidification due to the obvious ease of regenerating sulfuric acid from the sulfate byproduct. Even easier perhaps, the nitrate can be decomposed thermally to yield NO2 and O2 for making nitric acid or other uses. The usual routes of synthesis are not suitable for this of course as they defeat the purpose, either involving nitric acid itself or its derivatives, or resulting in the very highly stable sulfates I would like to avoid.

Here’s my idea. Electolysis with a copper anode in ammonium nitrate solution should normally break down the ammonium salt forming copper nitrate with liberation of ammonia at the cathode (possibly in a membrane or membraneless cell). However at least two further reactions would prevent this from yielding any copper nitrate; copper would be electroplated out from copper nitrate at the cathode, and ammonia at the cathode would complex with copper or copper compounds. To prevent this, I am considering an AC electrolysis process in a hot cell (near 100°C). The alternating current should disallow any copper from plating out (nitric acid would be retained), and as ammonia is eventually lost to evaporation from both electrodes during the cathodic period, the composition of the electrolyte should be driven toward copper nitrate.

The concerns I have is that the ammonia will be oxidized, resulting in overall loss of ammonia (which I would like to retain by bubbling into solution in an adjacent vessel), or complexed with copper resulting in contamination and losses.

In a DC cell with two copper electrodes, where copper is removed from the anode and deposited on the cathode, the electrolyte should still tend toward copper nitrate as ammonia should still be lost. The process would then be periodically reversed, but the AC electrolysis should take care of that automatically. The only difference really is the potential for ammonia to be oxidized. I don’t really know enough about the chemistry of ammonia/copper complexes to guess how this would affect the system, but upon boiling the electrolyte to dryness it should be easy to break down those impurities.

Because of the intended use of the final copper nitrate product, any remaining ammonium nitrate should be removed as well. Of course this is easily done by heating to decomposition, but the copper nitrate hydrate will decompose as well at those temperatures. I’m not really sure of another way to remove or destroy remaining ammonium nitrate or determine when there is no more present, except simply decomposing the lot and allowing ammonium nitrate to reform and solidify in a condenser.

The copper oxide left after the process can be recycled by smelting (even though that is an energy intensive process). Copper oxide could also be added to excess sulfuric acid used to dehydrate the HNO3, thereby converting it to copper sulfate for reclaimation. Though more tedious, the total process would allow production of fuming nitric acids without net loss or hydration of sulfuric acid or without requiring the difficult contact process to reclaim it.

[Edited on 6-3-2008 by kilowatt]

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Nah, nitrate will be reduced. It would be too easy if ammonia could be oxidized.

Ammonia is too soluble. I don’t see any reason why it should leave. Copper will be oxidized at the anode and taken into solution, while nitrate and H+ will be reduced at the cathode, losing NO3 and causing some alkalinity (the ammonium in solution will buffer the catholyte pH around 8-10). When the copper and alkali meet, Cu(OH)2 will form, which may dissolve in excess NH3 if you let it run for a really long time. If you use a membrane, this will not happen, but a buildup of NH3 in the catholyte will reduce conductivity (NH3 doesn’t ionize very much in H2O).

A better idea would be to fill the cathode chamber with a conductive, essentially unreactive alkali, like NaOH, placing your nitrate in the anode chamber. NaNO3 could be used. Na ions migrate to the cathode, so it becomes more basic (and more conductive, not at all a problem), while copper ions migrate from the anode into the anolyte. Note copper ions will want to migrate out as well; an ion-selective membrane would be great, but I suppose that’s a little too much to ask for! To minimize loss, you’ll have to change the anolyte when it’s only, say, 10% copper (since the amount of loss is proportional to the concentration, you should get logistic growth over time) and fumble with it from there.

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Try boiling copper oxide with a strong solution of ammonium nitrate, skip the electrolysis.

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Could use electrolysis for the synthesis of the copper oxide; I have. Copper anode and cathode, KNO3 electrolyte with good stirring. Copper hydroxide initially forms but the cell resistively heats to the point of decomposing it to the oxide.
Extended cell runs cause loss of nitrate by reduction to ammonia.

Tim’s Idea separating anolyte and catholyte would work, but would take a long time. I have found it very difficult to obtain synthetically useful currents in such cells, especially with improvised membranes.

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How would that work? Copper oxide is insoluble and both it and ammonium nitrate are far more stable than copper nitrate.

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NH4NO3 NH3 + HNO3 (far to the left, but concentrated solutions have pH 4 to 6)

Some of the ammoniua is sweep away with the steam, lthis also works with ammonium sulfate but gives NH4HSO4.

Try it, let maybe a tenth of the water boil away before replacing it, repeat several times. Works even better with carbonates.

However as the decomposition temperature of cupric nitrate is about the same as the melting point of ammonium nitrate, who’s own decomposition temperature isn’t much higher, it’s not likely to be useful preparatively.

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I know for a fact molten ammonium nitrate likes to decompose to H2O + N2O instead of the HNO3 + NH3 that I would prefer. Evidently this is not the same as in the boiling solution you describe?

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NH3 passing off as gas is the driving force. NH3 + CO2 boils off easier, hence the carbonate suggestion. CuO isn’t very basic, so it’s going to go slowly. Carbonates usually react faster (I would imagine the mildly acidic pH would help bring some into solution), and ammonium bicarbonate is a more favorable «leaving group».

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I had thought briefly about copper hydroxide but did not think the reaction would proceed. Are you suggesting copper carbonate and ammonium nitrate to remove ammonia and CO2 leaving copper nitrate?

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Well, there is no copper carbonate as such, just basic carbonates. They react faster than the oxide, especially oxide that’s been heated to a high temperature like copper scale.

Anything that drives the reaction balance in the desired direction is good. Boiling to help drive off ammonia, the steam leaving helps carry the ammonia away so boiling tends to be more effective than heating to just below boiling. Carbonates are better because the CO2 escapes the same way, so BaCO3 and MnCO3 will react fairly quickly. Strong bases, Ba(OH)2 for example, readily displace the weak base «NH4OH»; with strong bases you often can get away with a mildly warm solution and bubbling clean air through it to carry away the NH3.

Cu(OH)2 likely works faster than CuO, but you can’t heat it much or the hydroxide goes over to the oxide. And you were proposing to make nitric acid by heating cupric nitrate, which leaves the oxide so a method that cycles between oxide and nitrate seems simpler.

BTW
2Cu(NO3)2 5H2O => 2CuO(s) + 4NO2(g) + O2(g) + 5H2O ==> 4 HNO3 + 3 H2O
(ignoring how you get that to go to completion) limits the concentration of the nitric acid you’ll get. It is pretty strong, if you can get close to that limit you should be happy.

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Thanks for the help, looks like this may be workable.

Yeah I was going to use concentrated H2SO4 to dehydrate the HNO3 fully, but I was going to then react the left over sulfuric acid solution with the left over CuO to get CuSO4, which I was going to decompose to SO3 and CuO.

To go from CuO to (CuCO3*Cu(OH)2) is not a big deal really, it can be done with precipitation. Clearly going from CuO to 2Cu(NO3)2.5H2O and back is simpler though and I doubt going to the trouble of making the carbonate would save any time even if the reaction is much faster, because its production is itself time consuming.

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be very careful with Copper around ammonium nitrate, it can form an unstable salt (tetramine copper 2 nitrate) that can explode!

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Any idea how to make that not happen?

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Once that I needed some copper(II) nitrate whatever hydrate I prepared it by dissolving some copper wire in dillute HNO3 with stoichiometric H2O2 added (I was a bit stingy on HNO3). I remember that after I had evaporated and recrystalized it and was vacuum filtering it, I thought it would be good to wash the remains of water from the crystals by washing with acetone so it would dry rapidly. What a stupid thing to do! To my surprise I learned that copper(II) nitrate is well soluble in acetone.

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Here Ca(NO3)2 is hard to find as agricultural supplement but is easily made from NH4NO3 + lime..

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it`s still an undesirable product that present a potential Danger (not to mention legalities).

for this Not to be presented would be negligent!

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How about adding an excess of Cu metal to dilute (38%) nitric acid, filtering under suction to remove unreacted Cu and other insoluble junk, then evporate until crystals form? I’ve tried it; it works! Just beware of the noxious fumes of NO2- literature states that only concentrated nitric acid produces brown NO2 fumes; but from my experience, even dilute (38%) nitric acid produces NO2 when it reacts with Cu.

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actually I process All my scrap copper in this way, I save it all up in a jar and when I have a good 500g or so, I dissolve the lot in 38% Nitric with copper in excess.

then I convert it all to the Carbonate and store it in a jar.

from this point I can make any simple copper salt I like with ease.

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Dissolution of copper in dilute HNO₃ produces NO, a colorless gas that immediately oxidizes to NO₂ upon contact with atmospheric oxygen. More concentrated HNO₃ reduces mostly only to NO₂ during the oxidation of copper. By adding H₂O₂ you get a dissolution with nearly no gas formation until all the peroxide gets consumed. Also, as a consequence you need less HNO₃ to dissolve the same amount of copper (see the appropriate redox equations).

Atached (I have not read it, but I guess it is relevant even though ancient):
The Conditions of the Reaction between Copper and Nitric Acid
V. H. Veley
Proceedings of the Royal Society of London, 46 (1889) 216-222

[Edited on 7/3/2008 by Nicodem]

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I am trying to make nitric acid in the first place here from ammonium nitrate using a cyclic process, not the other way around. For me this is better than the acidifying an alkali nitrate because I do not have a huge supply of sulfuric acid to just be using up.

[Edited on 7-3-2008 by kilowatt]

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Why not just heat inorganic copper salts with a nitrate? Heating under a low flame (bunsen burner) a mixture of CuSO4. 5 H2O with KNO3 turns green and then forms brown nitrogen oxides, leaving a black residue.

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I am trying to find a route that uses ammonium nitrate specifically because that’s what I have most of, releases ammonia for collection, and does not produce any alkali metal salts of sulfuric acid (which are relatively difficult to recycle). The reaction of of CuSO4. 5 H2O with KNO3 would require first the reaction of KOH or NaOH with ammonium nitrate to produce said nitrate, and then the production of K2SO4 or NaSO4 by CuSO4 which ultimately uses up copious amounts of sulfuric acid by converting it to a relatively irrecoverable form.

I thought that’s what I tried to convey with my last post but here I am repeating myself.

I am trying to find overall reactions that simply produce nitric acid and ammonia from ammonium nitrate without using up anything else, not trade one acid for another. Copper salts probably aren’t the way to go after all (TACN can be produced), but I am looking at possible lead, aluminum, and magnesium routes involving electrolysis or amalgam electrolysis in other threads.

I think we all know how to produce Cu(NO3)2 or NO2/O2 for their own sake, and thanks for the additional route but this is a different process.

[Edited on 10-8-2008 by kilowatt]

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The mods should remove those posts. I will start another thread on this, since I haven’t seen this discussed.

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This thread is misnamed and meandered far afield from the propsed premise,extracting nitric acid
from ammonium nitrate. I don’t know if this can be made to work but it seems to me that a voltaic
cell can serve this purpose. The interfacing reservoir collects concentrated nitric acid and serves
for what is normally the connecting salt bridge.

What metal should be used to make a vessel to store concentrated nitric acid

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I usually make concentrated nitric acid by the same simple procedure as presented many times over, like here, except I use 30 cm Vigreux column for the second distillation.

My crude nitric acid (before fractionation) was 53%, the solution started boiling at 115 °C (239 °F) and the temperature at the still head was 110 °C (230 °F) instead of expected 100 °C (212 °F). So instead of removing mostly water, lots of acid has been removed as well.

Then I wait until the head temperature raises to about 118 °C (244.4 °F) and switch the flasks. If I wait longer, a good half of the volume in the boiling flask gets over before the head temperature finally reaches 120 °C (248 °F) which is almost the b.p. of nitric acid azeotrope. The fraction taken above 118 °C (244.4 °F) has a concentration of 58%, not much close to 68%.

I don’t understand why the Vigreux column provided worse fractionation than a simple distillation people often use to make 68% nitric acid?

I used aluminium foil insulation on the column but maybe it is not beneficial in this case? But even if the column was perfectly insulated, it would still at worst perform like a simple distillation, which is known to provide azeotropic or near-azeotropic nitric acid.

Should longer column fix the problem?

Since my fume hood is too small for an entire distillation apparatus, I use a hose attached to the vacuum adapter to take the nitrogen dioxide into the fume hood and then outside. Maybe just having the long hoe attached raises the pressure in the apparatus enough to spoil fractionation? I think this is improbable but I am out of ideas.

I will try titration to measure concentration more precisely, sine maybe I made a mistake in the density measurement (I use volumetric flask and a scale).

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First, it would be simpler to initially add only the quantity of water requied to produce (slightly greater than) azeotropic concentration in the original reaction,
this way the distillation is not actually required, just removal of the sulphate salt (and other impurities) by cooling and filtration,
A simple distillation would of course give a purer product.

As (drain unblocker) sulphuric acid is relatively cheap,
it can be used to de-hydrate dilute nitric acid to any concentration prior to distillation
===============================================
Here is a diagram of nitic acid vapour and liquid concentrations vs temperature
( taken from https://www.chemguide.co.uk/physical/phaseeqia/nonideal.html )

Looking at that diagram you can estimate that two ‘theoretical distillation plates’ should concentrate 53% to very nearly azeotropic,
the boiling pot/rbf is one ‘plate’ and a Vigreux column should easily provide the rest.
So in theory, from your 53%, you should have obtained near azeotropic nitric acid.

I like to use excess sulphuric acid and no additional water in the starting reaction,
on distillation this gives RFNA and some waste NO₂,
AFAIK RFNA is useful for most nitrations, and can be diluted to azeotropic if required.

[Edited on 9-4-2019 by Sulaiman]

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Drop the Vigreux (not literally) it’s useless.
It only forces you to use higher temperatures and decompose your product.

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Quote: Originally posted by Sulaiman

First, it would be simpler to initially add only the quantity of water requied to produce (slightly greater than) azeotropic concentration in the original reaction, this way the distillation is not actually required, just removal of the sulphate salt (and other impurities) by cooling and filtration, A simple distillation would of course give a purer product.

But how do you tell the reaction is complete? Have you done this filtration or is it just a hypothesis? I have not found this kind of work up elsewhere.

I thought the reaction starts at high temperatures so at least refluxing for several hours would be needed and if so, then it is as easy to perform distillation.

Quote: Originally posted by Sulaiman

I like to use excess sulphuric acid and no additional water in the starting reaction

This is known to cause significant loss of yield. Yes the nitric acid will be fuming and of higher concentration, but the total amount of HNO3 is lower. Diluting fuming acid itself produces lots of decomposition and fumes so this is why I like to add water beforehand, not after the fact.

I don’t need RNFA as I don’t do any nitrations. I need 68% specifically for dissolving metals like silver and making aqua regia. Higher concentrations don’t work for due to passivation of the metal and lower concentrations also work worse due to excess water and thus need for prolonged drying. The azeotropic one is used in most syntheses I like to reproduce, hence the need for 68%.

Quote: Originally posted by Sulaiman

Members more experienced than myse lf would perform nitric acid distillations at reduced pressure, hence reduced temperatures, as nitric acid is decomposed fairly quickly by higher temperatures.

Yes I can do vacuum distillation at 100 mbar. I have a chemically resistant pump so this is not an issue.

I didn’t know high temperature decomposes nitric acid. I always thought just that the light is the issue.

One problem with vacuum distillation is the shifted boiling points. The crude manometer on the pump can’t tell the exact pressure needed to estimate the boiling point using P-T nomograph and even if pressure is known, it is only an estimate. I have to observe the thermometer the whole time for any changes as this can be the new fraction coming.

Quote: Originally posted by Herr Haber

Drop the Vigreux (not literally) it’s useless. It only forces you to use higher temperatures and decompose your product.

But how do you fractionate then??

Should I use packed column with Rashig rings?

Yes there are some special tray-based fractionating columns but these are huge and very expensive. I know people make 68% nitric acid with only minimum equipment, in low tech amateur settings (e.g. Doug’s Lab on YT). So I know I am making a mistake somewhere and not sure where it is.

Maybe the higher temperature required for Vigreux column and the low performance of it actually leads to lower concentration than simple distillation?

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As for the density calculation, yes I checked the temperature and it was 28 °C. For the density of 1.34 g/mL, the table gives 58% for both 25 °C and 30 °C so I am confident about that concentration as it should fall between the two numbers. But I will do a titration anyway.

I am thinking about using a 40 cm packed column (filled with ceramic rings) and vacuum distillation to concentrate the acid.

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Can’t you get close to 68% nitric acid simply by boiling down the diluted acid? I think that once the concentration gets over 20%, that some acid would be lost through boiling. If set up for simple distillation these could be recovered, and additional concentrated acid could be recovered by re-boiling the distillate. It seems like a simple, if inefficient way of doing it.

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Quote: Originally posted by WGTR

Can’t you get close to 68% nitric acid simply by boiling down the diluted acid? I think that once the concentration gets over 20%, that some acid would be lost through boiling. If set up for simple distillation these could be recovered, and additional concentrated acid could be recovered by re-boiling the distillate. It seems like a simple, if inefficient way of doing it.

I want to use the acid for organic syntheses as well so I don’t want any salts in it, hence the need for two distillations in the original synthesis.

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Okay so I repeated the fractional distillation using vacuum and the separation was still poor.

The liquid started boiling at 66 °C (100 mbar vacuum) and the head temperature was 60 °C meaning there is probably a good separation.

The temperature if the liquid rose to about 72 °C and head temperature to 70-71 °C showing there is no longer a good separation.

I assumed the mixture is already azeotropic and stopped the distillation.

The amount of distillate was about 150 mL which is about the amount of water expected.

Unfortunately the concentration of acid was only 59%. So the distillation removed 10% of the volume, increasing the concentration by about 1% only.

It seems nitric acid cannot be efficiently concentrated with simple or fractional distillation unless very long column is used or some agent added (like sulfuric acid) to trap the water.

I think I will lower the amount of water in the initial synthesis to get near-azeotropic acid from start, even at the cost of lower yield, and then sacrifice 1/10th of the volume of the liquid to get to 68%.

The image shows the setup I used.

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What is it that you’re using to measure density? Density is very temperature sensitive. I would consider a titration to confirm your suspicions.

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Quote: Originally posted by UC235

What is it that you’re using to measure density? Density is very temperature sensitive. I would consider a titration to confirm your suspicions.

I used two methods, density calculation and titration. Both methods gave me 58% (by titration it was 57.97%) so I am confident about the result.

I used titration to measure acid concentrated by vacuum fractional distillation it was now 59.5%.

For titration, I use 10 g sample which slighly diluted with ddH2O and titrated with 20% NaOH from 50 mL burette (0.1 mL precision) and use few drops of nitrazine solution as neutralization indicator. The solution is magnetically stirred and the neutralization point is found by sharp color change which occurs within one drop of addition. I think it is precise to 1%

For density measurement, I use 250 mL volumetric flask (class A) and 0.01 g scale (calibrated).

I will research if small amount of sulfuric acid would help with concentration (I need 68% though, I have no use for fuming nitric acid) and I will probably try fractional distillation with 80 cm packed column. Hopefully this will finally help. Unfortunately, my packed column tends to overfill so maybe I will start with 40 cm one and see how it performs.

[Edited on 11-4-2019 by nimgoldman]

[Edited on 11-4-2019 by nimgoldman]

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I just measured the HNO3 concentration in the receiving flask and it was 29%.

I think the only viable way would be to add small amount of sulfuric acid to the mixture, as a dehydrating agent, and use simple distillation.

The question is how much sulfuric acid to use? 50% sulfuric acid has a boiling point of 120 °C which is just about the b.p. of azeotropic nitric acid (121 °C). However, as the water will boill off, the b.p. of the mixture will quickly rise and the azeotropic nitric acid will be the lowest boiling component. Clearly an excess of sulfuric acid will be needed, as well as re-distilling the mixture and using using fractionation.

I am a bit worried about sulfuric acid contamination though.

[Edited on 11-4-2019 by nimgoldman]

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There is a book in the Sciencemadness library entitled the “Absorption of Nitrous Gases”. Page 279 begins a chapter on the concentration of dilute nitric acid by various means, and provides some data that may be useful to you. I have this book, so if there are some pages that are difficult to read just let me know and I can scan them.

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From the book, boiling down an approximately 35% solution of nitric acid down to half its volume should give you a 65% concentration in the residue of the boiling flask. If you did a simple distillation on the diluted acid, then the first 2/3 or so could be your diluted fraction. The next fraction should be darn near 65-68%, in my understanding.

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I will try that.

Quote: Originally posted by WGTR

There is a book in the Sciencemadness library entitled the “Absorption of Nitrous Gases”. Page 279 begins a chapter on the concentration of dilute nitric acid by various means, and provides some data that may be useful to you. I have this book, so if there are some pages that are difficult to read just let me know and I can scan them.

Thanks. This is exactly what I looked for.

I will probably try vacuum distillation to get pure concentrated one and then concentrate the leftovers by adding sulfuric acid and distill them over.

[Edited on 12-4-2019 by nimgoldman]

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I decided to concentrate my 58% nitric acid to >68% by adding sulfric acid and distill under vacuum. If the resulting nitric acid will be >68%, I will dilute it afterwards.

There is one interesting graph in the book I don’t completely understand:

It seems that adding H₂SO₄ to my 57% HNO₃ (so that final concentration is 10% H₂SO₄) wil yield >85% HNO₃, right?

I don’t understand how long the distillation should run and what is the stopping condition.

There is one graph showing boiling point curves for various mixtures of nitric acid and sulfuric acid, which might give a hint about the stopping temperature:

From it, this seems that adding H₂SO₄ to 10% w/w concentration in my 57% nitric acid will raise boiling point to about 120 °C then slightly going up and then decrease to 110 °C as the nitric acid boils off from the mixture.

But this is problematic, since the dilute sulfuric acid has a lower b.p. than 120 °C and would distill over.

I really need to avoid contaminating nitric acid with sulfuric. I use nitric acid for things like dissolving silver and I really need pure nitrates, not sulfates.

One solution would be to add excess sulfuric acid (say 60% w.r.t. contained water in the solution), then its b.p. will always be well above the b.p. of nitric acid at any concentration. This will however produce fuming nitric acid, which I found difficult to dilute without too much decomposition and NO₂ production. it would probably have to be redistilled under vacuum.

I think any distillation carried out should be fractional, but this detail has not been mentioned in the book.

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You make me smile. Thank you

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Quote: Originally posted by nimgoldman

But this is problematic, since the dilute sulfuric acid has a lower b.p. than 120 °C and would distill over. I really need to avoid contaminating nitric acid with sulfuric. I use nitric acid for things like dissolving silver and I really need pure nitrates, not sulfates.

The amount of sulfuric acid coming over in a distillation should be almost nothing at 120°C. It’s some water that’s coming over, with the nitric acid.

I’m still not convinced that you need sulfuric acid at all anyway. It’s Friday. Let me see if I can run a small experiment today. If I can, I’ll post the results.

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Perhaps you should try to do everything the opposite you think makes sense. Retort still, conc. sulfuric acid, just enough heat for distillate, taking fractions when heat is increased.

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Quote: Originally posted by S.C. Wack

Perhaps you should try to do everything the opposite you think makes sense. Retort still, conc. sulfuric acid, just enough heat for distillate, taking fractions when heat is increased.

That’s a good idea. I like to use vacuum however, as this produces nice clearless nitric acid (without vacuum, I usually have lots of nitrogen dioxide produced).

I will probably attach a small fractionating column (just to be on the safe side) do vacuum distillation and collect fractions.

I just realized nitric acid is actually produced from sulfuric acid solution (in the most common synthesis) and there is no noticeable contamination.

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To concentration does not seem to work according to paper. They suggested as little as 10% sulfuric acid will allow producing fuming nitric acid from fairly dilute one, or I understood the graphs wrong.

Anyway, I took 300 mL of crude 47% nitric acid, added 30 mL of 98% sulfuric acid to it and distilled under vacuum (100 mbar) with a short Vigreux column. The temperature at still head slowly raised from 61 °C to 68 °C until the temperature slightly dropped and sulfuric acid vapours appeared in the boiling flask at which point I stopped the distillation.

There was about 270 mL of distillate with a density of 1.3 g/mL corresponding to 47% concentration.

So the addition 10% of the volume in sulfuric acid didn’t help in any way. Should I use more sulfuric acid?

Now I am thinking about producing fuming nitric acid by adding excess of sulfuric, then concentrating the near-concentrated nitric acid I have with the fuming one.

I will probably perform the first distillation (with sulfuric acid) without vacuum to monitor the actual temperature so I can compare it to tables, then redistill it under vacuum to remove nitric dioxide contamination.

[Edited on 15-5-2019 by nimgoldman]

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Why don’t you try producing fuming nitric acid directly from sulphuric acid and potassium nitrate, then you dilute it?

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Quote: Originally posted by nimgoldman

To concentration does not seem to work according to paper. They suggested as little as 10% sulfuric acid will allow producing fuming nitric acid from fairly dilute one, or I understood the graphs wrong. Anyway, I took 300 mL of crude 47% nitric acid, added 30 mL of 98% sulfuric acid to it and distilled under vacuum (100 mbar) with a short Vigreux column. The temperature at still head slowly raised from 61 °C to 68 °C until the temperature slightly dropped and sulfuric acid vapours appeared in the boiling flask at which point I stopped the distillation. There was about 270 mL of distillate with a density of 1.3 g/mL corresponding to 47% concentration. So the addition 10% of the volume in sulfuric acid didn’t help in any way. Should I use more sulfuric acid? [Edited on 15-5-2019 by nimgoldman]

The curves you posted are pretty clear: here you started with a 50% nitric acid combined with 10% sulphuric acid solution. If you look at the graph, the rightmost curve which represents 10% sulphuric acid, you will see that you’re not expected to get a much higher concentration of nitric acid in the fumes (you get even LESS). If you double the sulphuric acid concentration (20%) however, from a 50% nitric acid solution you should be able to boil off a 80% solution, at a boiling point of around 120 °C, as the second graph tells you. And 30% sulphuric acid should give you almost pure nitric acid (look up the 50% nitric acid vertical line).

At least that’s the way I interpret those curves.

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https://www.sciencemadness.org/whisper/viewthread.php?tid=13. This method is my favorite due to its simplicity. I have done it twice and both times I got acid with a density that corresponded well to 68%. I only used a molar equivalent of ammonium nitrate instead of KNO3 and then added a few drops of 50% H2O2 to my HNO3 to remove any residual nitrogen oxides that remained dissolved.

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I’ve ran the experiment with 900 mL of 47% HNO3. It contains about 595 g of water so I added about half of that volume in sulfuric acid, i.e. 300 mL.

The acid produced was only 58%, not over 80% like the curves in graph suggested.

I then realized most syntheses use as much sulfuric acid as there is water, even more and one synthesis of anhydrous nitric acid from conc. nitric acid requires five times the volume in sulfuric acid as there is nitric.

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