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Author: Subject: Nitric Acid Synthesis
hissingnoise
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[*] posted on 22-12-2009 at 09:15


Well, OK, but N2O5's main hazards are described thus; (strong oxidiser; forms strong acid in contact with water.).
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[*] posted on 23-12-2009 at 11:02


I took HS chem in 2007-2009, we had a separate course for labwork.
All the basic extraction, titration and chromatography but also we oxidized stuff with nitric acid.
During the last two classes every group manufactured a small ~10g batch of black powder.
Next year the next group did a "bengal fire" demo with sodium chlorate but we actually got to make the black powder and burn it too. We also used some old thermite leftover from some old class.

Our chem book also had a small two pages of text and one page of exercises about explosives. Some highlights of the exercises "calculate the gas output of this ammonium nitrate and diesel oil mixture", "write out and balance the synthesis of nitroglycerin" and there was also one about tnt synthesis.
Most interesting I say. Needles to say that this was not in any english speaking nation, but still in Europe.

But now on the subject.

I'm probably going to try the "DCM method" for producing some nitric.
Also, if nitric oxide contamination is not a problem how about NaHSO4 (l) + xNO3 (s) --> HNO3?
NaHSO4 melts at ~58*C, below the boiling point of nitric acid. And molten NaHSO4 is said to behave like concentrated sulfuric acid, which is hard for me to get.

[Edited on 23-12-2009 by iHME]




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[*] posted on 23-12-2009 at 11:04


How is this relevant ?




What a fine day for chemistry this is.
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[*] posted on 23-12-2009 at 11:09


'Tis the season to be, er, tight?
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[*] posted on 23-12-2009 at 13:57


If I have read right there was someone in this thread complaining about the lack of lab work in HS chem.
I used that as a excuse to brag how much different it was for me, with the disguise of it being about how things can be different even if one lives in a paranoicratic western country. And I did post some stuff relevant to the thread's subject too.

Anyways, once again back to the topic. How 'bout that idea about replacing molten NaHSO4 for concentrated H2SO4?
It "should" work, but it would produce nitric acid contaminated with nitric oxides and the temperature might promote excessive destruction of the formed acid. But it should be cheap to manufacture and the reagents would be easy to get, but people might look at you funny if you buy pool pH- when it is -15*C outside. ;)




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[*] posted on 23-12-2009 at 14:17


You edited your post afterwards and you didnt have the " now on subject part "
It was edited in later.




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[*] posted on 23-12-2009 at 14:22


I hope Iceland will accommodate all of us, because we're all coming, iHME.
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[*] posted on 23-3-2010 at 14:15


Last weekend I produced aqueous nitric acid by the direct acidification of ammonium nitrate by an equimolar amount of sulfuric acid. If I remember correctly is a process which had been proposed here before, but no one was quite sure how well it would work. I decided to find out, and it does. The next step will be to determine if fuming nitric acid can be produced in a similar process under vacuum distillation.

I used an equal amount of water by volume to the concentrated sulfuric acid, which the acid was added into prior to adding the ammonium nitrate. This would be equivalent to about 60% H2SO4 by weight or 90%w/v, which should be easily obtainable by concentrating battery acid if someone should want to go that route. The negative heat of solution of the ammonium nitrate helped to absorb the positive heat of solution of the sulfuric acid upon addition, although both additions were done quite gradually and I have not calculated the total enthalpy.

I distilled it until the reaction liquor in the flask became frothy with a volume about 1.5 times the original volume, yielding a little over 500mL of acid per liter of raw reaction liquor. The main portion of the distillate came over at 113°C (evidently the boiling point of the azeotrope at my altitude) after a brief fraction at 94°C (the boiling point of water at my altitude). I did not drain the water prior to the main fraction, and the total product had a specific gravity of 1.36. There were no signs of ammonia (which I feared could be produced as decomposition product) at any stage of the distillation, which would have given a smoky appearance to the vapor. Once the distillation stabilized, the acid came over with a slight yellow tint and some NO2 color was visible inside the condenser. This was cleared up by bubbling pure oxygen through the obtained acid for a few minutes, although air could have been used with reduced efficiency.

I am in the process of dissolving the spent reaction liquor, which is a solid mass upon cooling and consists primarily of ammonium bisulfate with a small excess of sulfuric and nitric acids. Once this is done I intend to the react the solution with an excess of either zinc oxide or copper oxide, which should proceed (probably slowly) according to the reaction CuO + NH4HSO4 --> CuSO4 + NH3 + H2O. The resulting sulfate will then be suitable for thermal decomposition to yield SO3 after preliminary heating has removed all the left over nitrate from the system in the form of NO2 by decomposition of copper nitrate. Zinc sulfate will give a lower yield of SO3, giving instead a mixture of SO3 and SO2. The latter could be used to produce a bisulfite for unrelated uses, or could be catalytically oxidized to yield SO3. I am under the impression, however, that CuSO4 decomposes pretty selectively to SO3 and CuO; correct me if I am wrong. Zinc oxide/sulfate would certainly be a cheaper way to go in terms of how much material is tied up in this cyclic process, but the decomposition is not as favorable. Are there any other oxides that would be good candidates for this?




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[*] posted on 16-4-2010 at 20:48


On the note of Nitric acid, I landed a job in Nitric acid plant and was looking at the flow/density meter and regression for acid concentration from density and temperature.

I found tables but no luck on the equation, does anyone have strong enough interest in HNO3 to have looked into this?

The current regression we have swings quite erratically for the past few weeks, up and down by 2% compared to the lab assay.

Quote: Originally posted by kilowatt  
Last weekend I produced aqueous nitric acid by the direct acidification of ammonium nitrate by an equimolar amount of sulfuric acid. If I remember correctly is a process which had been proposed here before, but no one was quite sure how well it would work. I decided to find out, and it does. The next step will be to determine if fuming nitric acid can be produced in a similar process under vacuum distillation.

I used an equal amount of water by volume to the concentrated sulfuric acid, which the acid was added into prior to adding the ammonium nitrate. This would be equivalent to about 60% H2SO4 by weight or 90%w/v, which should be easily obtainable by concentrating battery acid if someone should want to go that route. The negative heat of solution of the ammonium nitrate helped to absorb the positive heat of solution of the sulfuric acid upon addition, although both additions were done quite gradually and I have not calculated the total enthalpy.

I distilled it until the reaction liquor in the flask became frothy with a volume about 1.5 times the original volume, yielding a little over 500mL of acid per liter of raw reaction liquor. The main portion of the distillate came over at 113°C (evidently the boiling point of the azeotrope at my altitude) after a brief fraction at 94°C (the boiling point of water at my altitude). I did not drain the water prior to the main fraction, and the total product had a specific gravity of 1.36. There were no signs of ammonia (which I feared could be produced as decomposition product) at any stage of the distillation, which would have given a smoky appearance to the vapor. Once the distillation stabilized, the acid came over with a slight yellow tint and some NO2 color was visible inside the condenser. This was cleared up by bubbling pure oxygen through the obtained acid for a few minutes, although air could have been used with reduced efficiency.

I am in the process of dissolving the spent reaction liquor, which is a solid mass upon cooling and consists primarily of ammonium bisulfate with a small excess of sulfuric and nitric acids. Once this is done I intend to the react the solution with an excess of either zinc oxide or copper oxide, which should proceed (probably slowly) according to the reaction CuO + NH4HSO4 --> CuSO4 + NH3 + H2O. The resulting sulfate will then be suitable for thermal decomposition to yield SO3 after preliminary heating has removed all the left over nitrate from the system in the form of NO2 by decomposition of copper nitrate. Zinc sulfate will give a lower yield of SO3, giving instead a mixture of SO3 and SO2. The latter could be used to produce a bisulfite for unrelated uses, or could be catalytically oxidized to yield SO3. I am under the impression, however, that CuSO4 decomposes pretty selectively to SO3 and CuO; correct me if I am wrong. Zinc oxide/sulfate would certainly be a cheaper way to go in terms of how much material is tied up in this cyclic process, but the decomposition is not as favorable. Are there any other oxides that would be good candidates for this?
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[*] posted on 17-4-2010 at 23:05


Along the way I wondered, how a used old car-catalyst might work for synthesizing Nitric acid, by maybe burning NH3 or even for synthesizing NH3 from synthesis-gas (blowing steam through hot coal ...) :
==> If possible, this could "democratize" the HNO3-Production ... :D

[Edited on 18-4-2010 by chief]
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[*] posted on 18-4-2010 at 09:56


Cat converter in car is designed to actually do otherwise, convert NOx to N2, can do it 2 ways:
1. Selectively, use NH3 (from urea) to reduce NOx to N2
2. Non Selectively, use excess fuel (hydrocarbon), again to reduce NOx to N2.



Quote: Originally posted by chief  
Along the way I wondered, how a used old car-catalyst might work for synthesizing Nitric acid, by maybe burning NH3 or even for synthesizing NH3 from synthesis-gas (blowing steam through hot coal ...) :
==> If possible, this could "democratize" the HNO3-Production ... :D

[Edited on 18-4-2010 by chief]
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[*] posted on 19-4-2010 at 03:21


I thought of that; but at the end the car-cat just only contains Platinum, Palladium and some Rhodium, on ceramic substrate ..
==> So it should, when driven within the right parameters, do anyhing that such a cat would usually do ...

I'm a catalysis-greenhorn ...; maybe I heared somewhere about the "reducing" properties of those platinum-metals ... when they get into the human body ...

Also the standard-ctalyst for making HNO3 wa V2O5 or something ... ???

Anyhow: Was just an idea ...; if it would work, then probably it would work great, since the catalyst would be well manufactured ...
=================
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[*] posted on 23-4-2010 at 19:39


No, you were somewhat right with the concept.
The cat for making NO (and thus HNO3) is also made from Platinum Rhodium Palladium (there are a few different alloys available in the market), however the difference is in the design and operating temperature.
Usually the cat. for oxidizing NH3 -> NO is in the form of pure metallic gauze, relatively thin to control the contact time (because extended contact can actually reverse back the reaction to N2). The op. temp is also higher around 1500-1600 F vs. 1100-1200F for NOx abatement; at this higher T, the car auto converter might actually start sintering.

V2O5 is usually used in contact process for making H2SO4.


Quote: Originally posted by chief  
I thought of that; but at the end the car-cat just only contains Platinum, Palladium and some Rhodium, on ceramic substrate ..
==> So it should, when driven within the right parameters, do anyhing that such a cat would usually do ...

I'm a catalysis-greenhorn ...; maybe I heared somewhere about the "reducing" properties of those platinum-metals ... when they get into the human body ...

Also the standard-ctalyst for making HNO3 wa V2O5 or something ... ???

Anyhow: Was just an idea ...; if it would work, then probably it would work great, since the catalyst would be well manufactured ...
=================
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[*] posted on 23-4-2010 at 20:01


that would make sence arsen otherwise our
cars would be giving off acid rain.

I personaly can not see a way to make pure alloy gauze

especially one that is so noble.

it would have to be brought.

maby the black that is given under high temp and hydrogen
from the ammonium chloride salts of the nobles might do it.

I know this makes a kind of sponge though that would
probably be too thick and give the result that you are
describing.

amazing stuff the platinum group.


[Edited on 24-4-2010 by Ephoton]




e3500 console login: root
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[*] posted on 23-4-2010 at 20:09


Quote: Originally posted by chief  
Along the way I wondered, how a used old car-catalyst might work for synthesizing Nitric acid, by maybe burning NH3 or even for synthesizing NH3 from synthesis-gas (blowing steam through hot coal ...) :


The platinum group metals are rather poor at forming NH3 from H2 and N2, the preferred catalysts are based on iron doped with alkali metal oxides. These catalysts are poisoned by carbon monoxide, considerable effort goes into removing CO from the syngas stream; first by making more H2 via the water shift reaction CO + H2O <=> CO2 + H2, then by scrubbing CO2 out of the H2, and finally by reducing traces of the carbon oxides to methane and water. the CH4 being fairly inert towards the catalyst.

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[*] posted on 23-4-2010 at 22:57


The main reason why a car cat is scrap is that they do get contaminated in service, though now not so much as when leaded petrol was on the forecourt!
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[*] posted on 24-4-2010 at 00:07


Quote: Originally posted by kilowatt  
...

I am in the process of dissolving the spent reaction liquor, which is a solid mass upon cooling and consists primarily of ammonium bisulfate with a small excess of sulfuric and nitric acids. ... Are there any other oxides that would be good candidates for this?


What you need is a copy of High Temperature Properties and Thermal Decomposition of Inorganic Salts with Oxyanions by Kurt H. Stern. You may be able to see parts of it in Google Books, search for sulfate thermal decomposition

The PDF at http://cpercla.org/pdfs/pubs/tu/copper.pdf may be of interest.

From a quick bit of research it appears that in general the decomposition goes as

M<sub>x</sub>(SO4)<sub>y</sub> => "M<sub>x</sub>O<sub>y</sub>" + y SO3

SO3 => SO2 + 1/2O<sub>2</sub>

The lower the decomposition temperature of the sulfate the greater the SO3 fraction. CuSO4 gives a lot to mostly SO2.

Possibly increasing the O2 concentration would help suppress the breakup of SO3, as might quick removal and chilling of the SO3. I've been told that the portable oxygen "generators" used for medical patients are available used for decent prices, functional but no longer of medical grade. Blow very hot oxygen enriched air through the sulfate, then quickly quench the gases by passed through H2SO4 that is kept relatively cool.

Ferrous sulfate may work, as it was used long ago in making H2SO4. See http://www3.interscience.wiley.com/journal/114221358/abstrac...

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[*] posted on 25-4-2010 at 13:27


The oxygen concentrator/generators are usually retired for financial/accounting reasons, so the downgrading to non-medical use is paperwork only. I've measured approximately 93-95% O2 (the rest presumably Ar) at 1/2 of full flow. The purity of the output drops drastically as full output is approached. For glass, I use 5lpm units at 3.5-4lpm.

The last 5% Ar does lower flame temperature quite a bit. :mad:
Some day I'll look at the economics of putting together an O2/Ar separator - the ones I read about use synthetic zeolites like the N2/O2 separators but in a somewhat different configuration.
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[*] posted on 27-4-2010 at 05:36


Quote: Originally posted by densest  
The last 5% Ar does lower flame temperature quite a bit. :mad:
Some day I'll look at the economics of putting together an O2/Ar separator - the ones I read about use synthetic zeolites like the N2/O2 separators but in a somewhat different configuration.
The engineering handbook I read on pressure swing adsorption systems (what these concentrator boxes have inside them) didn't mention any material that did O2/Ar separation. Perhaps I missed it, or perhaps it's too new. It did say, specifically, that the zeolites to do N2/O2 separation did not distinguish between O2 and Ar, at least not those available at the time of publication. When you say "different configuration", it may be that it's not even a PSA system. Can you elaborate?

One useful thing I did learn from that book is that there's a different zeolite material that preferentially adsorbs N2 rather than O2/Ar; this can be used to generate rather pure N2 gas which can substitute for bottled N2 in most circumstances. Also, packing the adsorption columns with silica gel yields dry air that's dry in the ppm range, which seems like an excellent adjunct for someone whose desiccator is their rate-limiting piece of equipment.
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[*] posted on 27-4-2010 at 08:38


The above link ( http://cpercla.org/pdfs/pubs/tu/copper.pdf ) contains a maybe interesting way to produce H2SO4:

Apparently CuO + SO2 gives CuSO4 ... which could then be electrolyzed ...

If electrolyzed with too much anode-current/cm^2 this would give fine copper-powder, which easily could be oxidized to CuO again ....

The SO2 could stem from the thermal decomposition of maybe CaSO4 somehow ?
================

Making a batch of H2SO4 each other day ... ... :D

[Edited on 27-4-2010 by chief]
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[*] posted on 27-4-2010 at 08:55


Why not just buy CuSO4 - or the fungicide Bordeau Mixture which contains it?

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[*] posted on 27-4-2010 at 11:01


Quote: Originally posted by Alain123  
I know this has been posted before, and I have done a search on it.

[snip]



Dangerous Acids Made Safely by Home Chemist
Popular Science July 1934
http://tinyurl.com/3yt6s6n

Long gone are the days of the Home Chemist in any magazine.

I remember make nitric acid using rubber stoppers (good
for two uses) and rubber surgical tubing (good for one use —
on a good day.)

Be careful of the PSM Home Chemist method of making
synthetic rubber. (May 1945) Heating ethylene chloride
is best done out of doors!
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[*] posted on 27-4-2010 at 17:47


Making metal gauze in your garage is quite a challenging task, I think.

Metal sponge wouldn't work too well because at this temperature they might start sintering, and the reaction is hard to control too (would be way too violent), unless the stream is diluted down or the metal is supported on inert substrate.

Quote: Originally posted by Ephoton  
that would make sence arsen otherwise our
cars would be giving off acid rain.

I personaly can not see a way to make pure alloy gauze

especially one that is so noble.

it would have to be brought.

maby the black that is given under high temp and hydrogen
from the ammonium chloride salts of the nobles might do it.

I know this makes a kind of sponge though that would
probably be too thick and give the result that you are
describing.

amazing stuff the platinum group.


[Edited on 24-4-2010 by Ephoton]
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[*] posted on 28-4-2010 at 01:35


Quote: Originally posted by hissingnoise  
Why not just buy CuSO4 - or the fungicide Bordeau Mixture which contains it?



Because then you pay the copper: _Quite_ expensive, when compared to obtaining the SO2 from plaster ...

Anyhow no such fungicide in Germany ...

Plaster: Te bag (25kg) costs maybe 5 $ or less ...; thats a reasonable price for some raw-material ...
==> The CuO-CuSO4-cycle would just do the oxidation of the SO2 ==> SO3 ..., replacing more difficult alternatives ...

I bet the liter conc. H2SO4 could be had for 1 $ or less, this way ...
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[*] posted on 28-4-2010 at 04:10


Quote:

I bet the liter conc. H2SO4 could be had for 1 $ or less, this way ...

But when you factor in the time and equipment needed it'll end up costing more than the 14 Euro I pay for 96% H2SO4 draincleaner. . .
I'm surprised that H2SO4 is so difficult to get in Germany.

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