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Author: Subject: help make white fuming nitric acid.
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[*] posted on 28-5-2010 at 12:05


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



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[*] posted on 28-5-2010 at 12:45


instead of make a decomposition of your fuming nitric acid using high temperature you can use batter method:
put your nitric acid under vacuum(10-30 min at 200 mmHg or 27 kPa)and the dissolved nitrogen oxides will be removed, its will limit the decomposition of your wfna ... because its will be 20-25 c instead of 60-50 c, i think 50-60 Celsius are to high for wfna ... its just make terrible decomposition.
i guess that make a hno3/h2so4 distillation to obtain wfna ... and then use vacuum to remove nitrogen oxides will give the best results for the home Chemist.
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[*] posted on 28-5-2010 at 13:27


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

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

P.S.

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

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

WFNA decompose around 83 °C. Study harder.
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[*] posted on 28-5-2010 at 14:21


vacuum distillation is not practical because its lower the liquefied point of nitric acid ... so the hno3 need water which are more cold then ice-water mixture ....
i have vacuum pump .... how can i run get colder solution in condenser ... which allow me to use my vacuum ability?!
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[*] posted on 30-5-2010 at 22:36


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

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

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




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[*] posted on 31-5-2010 at 12:48
RFNA


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

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

What la book fails to note is that .... nothing happens for
a couple of minutes during which time you add more
formaldehyde... when suddenly you have in you Erlenmeyer
flask... a fuming nitric acid volcano, and a heck of a mess.
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[*] posted on 20-9-2011 at 11:37


By bubbling nitrogen dioxide, along with dry air, into azeotropic (68.5% concentrated) nitric acid, the concentration can be increased up to 77%.
http://www.patentgenius.com/patent/4064221.html
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[*] posted on 21-9-2011 at 12:01


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

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

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

subsequently compressing the gases;

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

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

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

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

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

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

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

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

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

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

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

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

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


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




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[*] posted on 21-9-2011 at 14:06


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

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

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

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

[Edited on 21-9-2011 by AndersHoveland]
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[*] posted on 21-9-2011 at 14:51
Reaction between NO2 and HCl to make HNO3


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

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

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

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

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

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

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


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

wiley:
11-26-2010

hey there Anders,

would like your opinion on something.

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

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

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

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

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

Anyways, what do you think?

Anders Hoveland:
Complete Answer to your Idea:

"wiley]hey there Anders,

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

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

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

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

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

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

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

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

wiley:
11-27-2010

Hey Anders,

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

Btw, you said:

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

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

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

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

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

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

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

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

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

wiley:
11-29-2010

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

In the description he says:

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

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

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

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

Anders Hoveland:

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

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

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

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

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

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

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

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

wiley:
11-30-2010

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

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

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

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

Anders Hoveland:
12-03-2010

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

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

Anders Hoveland:
12-03-2010

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

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

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

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

But do whatever is easiest and safest

Anders Hoveland:
12-04-2010

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

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

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

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

wiley:
12-06-2010

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

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

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

Anders Hoveland:
12-06-2010

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

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

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

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


[Edited on 21-9-2011 by AndersHoveland]
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quicksilver
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[*] posted on 22-9-2011 at 06:02


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

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

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




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[*] posted on 28-9-2011 at 08:33


Here, this will help you> http://www.youtube.com/watch?v=TpWwBxsyJok&list=FLjU9mst...
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[*] posted on 29-9-2011 at 13:43


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

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

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

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

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

[Edited on 29-9-2011 by AndersHoveland]




I'm not saying let's go kill all the stupid people...I'm just saying lets remove all the warning labels and let the problem sort itself out.
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Steve_hi
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[*] posted on 29-9-2011 at 14:41


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

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

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

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

upload 1.jpg - 61kB upload3.jpg - 55kB upload2.jpg - 35kB
The 3rd picture you can see the green where I put a piece of copper wire i had to drop a few drops of water to get it to react.
there was a lot of redish brown smoke and it was bubbling it was quite a stron eaction so I think the concentration was quite high.
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[*] posted on 29-9-2011 at 15:04


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

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

[Edited on 29-9-2011 by quicksilver]




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[*] posted on 29-9-2011 at 15:17


Thanks quicksilver ill ry tilting it downwards on my next batch also thanks for the info on the kimax i intend to buy a 24/40 distilation unit i'll makes sure its kimax. Its quite expensive not knowing what to buy and then end up buying twice.
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[*] posted on 29-9-2011 at 15:20


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

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

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[*] posted on 29-9-2011 at 16:51
Nitric acid storage


I was just in my lab making some FeClII and checked in on my nitric acid which I had put in a little glass jar i got at the dollar store which has a rubber seal. The rubber has melted away and the jar is all dirty now I gues the nitric acid ate the rubber away.
I was wondering if the plastic bottle which contained the sulfuric acid would be capable of storing the nitric acid.
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[*] posted on 29-9-2011 at 21:22


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

If you know any better correcet me.
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[*] posted on 30-9-2011 at 01:05
HDPE containers for Acid


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

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

Chemical_Reference_Summary-small[1].jpg - 144kB
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[*] posted on 30-9-2011 at 01:38


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

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

Well, good luck!
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[*] posted on 30-9-2011 at 06:12


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



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

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

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

[Edited on 30-9-2011 by quicksilver]




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[*] posted on 1-10-2011 at 14:35


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

Does this mean i can do a distillation using HCl and get nitric acid as well as ammonium chloride because HCl is much cheaper than sulphuric acid?
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[*] posted on 1-10-2011 at 22:26


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

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


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

Just use google...
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[*] posted on 1-10-2011 at 23:14


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

Does this mean i can do a distillation using HCl and get nitric acid as well as ammonium chloride ?


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

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

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

[Edited on 2-10-2011 by AndersHoveland]




I'm not saying let's go kill all the stupid people...I'm just saying lets remove all the warning labels and let the problem sort itself out.
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