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Author: Subject: How can I make niric acid?
tinker Terry
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[*] posted on 17-5-2013 at 22:46
How can I make niric acid?


I need nitric acid for a project I am working on and cannot find a local supplier.:( is there a reasonably simple way to make it at home with minimum expense and rudimentary equipment?



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[*] posted on 17-5-2013 at 23:03


You really want nitric acid? That is some of the nastiest stuff I've ever encountered.

Even with a proper distillery kit, it is still a real big PITA. Even if you successfully create nitric acid, even storing it is a big problem as well. You have to store it in an all glass container. Nitric will eat rubber and most things without any trouble.

BTW, you are in the same boat as I. Local suppliers will not give it to you because it is a chemical listed in the explosives act of Canada. What do you need it for? There are many good uses for nitric acid but sometimes you can use alternatives.

Anyways, see nurdrage's video regarding nitric acid if you really still want to make it: http://www.youtube.com/watch?v=2yE7v4wkuZU

If distillery is your choice of synthesis, another good video is this: http://www.youtube.com/watch?v=nRKAP7v3cv4


[Edited on 18-5-2013 by binaryclock]

[Edited on 18-5-2013 by binaryclock]




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weiming1998
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[*] posted on 17-5-2013 at 23:07


Here is a step-by step guide to find a simple way of making nitric acid:
1, Type www.google.com into the address bar.
2, Type "how to make nitric acid" into the search engine. Add "sciencemadness" at the end if you want to search this website.
3, ???
4, PROFIT

In other words, there is something called a search engine. Use it.

[Edited on 18-5-2013 by weiming1998]
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[*] posted on 18-5-2013 at 03:47


If you lived nearer to me, I would have sold you some. Nitric is not that bad at 70% conc. Stored in a polypropylene lined/caped bottle, nothing is corroded/leak. He is better to fallow magpie procedure in the prepublication sub forum. Else he will get acid too concentrated and contaminated with NOx. And then, he still need to find KNO3.



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[*] posted on 18-5-2013 at 05:26


Nitric acid can be purchased in Canada from specialty shops dedicated for plant growing. Search for hydroponics. It was 24$ for 500 mL or 1 L I think. It is very expensive. In some countries, it is 1 to 5$.
You can also buy a 25 Kg bag of Ca(NO)2 from such places.
I have produced some nitric acid acid using a high voltage circuit from a TV and air. The NO2 dissolves in water and makes HNO3 and NO. The NO reacts with O2 to make NO2 again. I made 0.93 M HNO3 this way, impure.
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[*] posted on 31-5-2013 at 05:19


I have not entirely performed this procedure myself, but you may consider and test it yourself for comparative purposes of cost and convenience for preparing Nitric acid if you do not have a nitrate or nitrite salt. Generally speaking, starting with a nitrate is a lower cost and more direct approach.

My suggested synthesis without a nitrate is as follows: first, add a copper source (copper plated pennies, etc.) in excess to a container along with an excess of aqueous ammonia (dilute or concentrated). Let stand for awhile and then add H2O2 (even dilute will work). Having performed this part, I can tell you that an obvious gas formation occurs on the surface of the copper and the solution temperature rises. A dark blue colored complex is formed and more gas (including N2) is evolved on standing, so do not seal your reaction vessel. A large sudden rapid spontaneous gas evolution is also possible (take precautions, see below).

Next step (untried by myself but confirmed) upon formation of the dark blue complex (this should occur relatively fast, but does vary depending on the concentration of your NH3 and H2O2) react with either H2SO4 and FeSO4 to produce NO (and NO2 in the presence of air). Let the NO2 dissolve in water or dilute H2O2 and further treat with air to form pure HNO3. One may be able to substitute NaI and Acetic acid with mild heating to 55 C for H2SO4 and FeSO4 to produce the Nitrous oxide (see "Effects of reducing reagents and temperature on conversion of nitrite and nitrate to nitric oxide and detection of NO by chemiluminescence", by Fan Yang, Eric Troncy, Martin Francœur, Bernard Vinet, Patrick Vinay, Guy Czaika and Gilbert Blaise, where the full paper is available free at http://www.clinchem.org/content/43/4/657.full although NaI/ Acetic acid is apparently not effective on liberating NO from any formed nitrates, only nitrites).

Chemistry source, see http://www.sciencemadness.org/talk/post.php?action=reply&... , to quote:

Quote: Originally posted by AJKOER  
Here is a more recent 1962 study on NH4NO2 formation examining various underlying theoretical models of the reaction called "Kinetics and Mechanism of Copper Dissolution In Aqueous Ammonia" fully available after signing on to ones Facebook account at http://www.academia.edu/292096/Kinetics_and_Mechanism_of_Cop... the author cites a rate for Cu dissolution as a function of available O2 and NH3.

Some of the underlying reactions cited by the authors include:

2 Cu + 4 NH3 + 1/2 O2 + H2O --> 2 [Cu(NH3)2]OH

2 [Cu(NH3)2]OH + 4 NH3 (aq) + 1/2 O2 + H2O --> 2 [Cu(NH3)4](OH)2

Cu + [Cu(NH3)4](OH)2 <---> 2 [Cu(NH3)2]OH

And, with respect to this thread, an important side reaction:

2 NH3 (aq) + 3 O2 + [Cu(NH3)4](OH)2 --> [Cu(NH3)4](NO2)2 + 4 H2O

Now, I actually performed the above reaction replacing atmospheric oxygen with some dilute H2O2 to speed things up. To my surprise, Copper pennies (my Cu source) became readily covered with O2 in agreement with a cathodic reduction reaction of oxygen at the copper's surface per the author's electrochemical dissolution model. The reaction is also apparently exothermic as the solutions became warmer. Within an hour, a dark blue was apparent. In 8 hours, a different lighter shade of blue was apparent that is characteristic of the usual cupric salts. Expected products could include tetraamminediaquacopper(II) dihydroxide, [Cu(NH3)4(H2O)2](OH)2, as well a monohydroxide, tetraamminecopper(II) nitrite and also the nitrate. The important side reactions forms NH4NO2, which was somewhat apparent by more excess gas formation than I suspected (do not used a sealed vessel) with the formation of both O2 and N2 (via a nitrite decomposition reaction).

Caution: The presence of Copper Ammonium nitrite and/or Ammonium nitrite may present a potential spontaneous nitrogen gas decomposition issue, which are more likely in slightly acidic or concentrated solutions. I would also be concerned on heating an acidified form of the solution just prepared due to known stability issues with hot aqueous NH4NO3 in the presence of metallic impurities (including Copper, Tin and Nickel see http://www.google.com/url?sa=t&rct=j&q=ammonium%20ni... ).
--------------------------------------------------

Here is a less authoritative 2011 study ("Copper-Mediated Non-Enzymatic Formation of Nitrite from Ammonia and Hydrogen peroxide at Alkaline pH" ) that is pertinent relating to nitrite formation noted above (please see http://www.google.com/url?sa=t&rct=j&q=reaction%20of%20nh3%2Ch2o2%20and%20cu&source=web&cd=4&ved=0CDwQFjAD&url=http%3A%2F%2Fsp hinxsai.com%2Fvol3.no2%2Fchem%2Fchempdf%2FCT%3D23(646-656)AJ11.pdf&ei=iS-mUfCNN4nr0gGYw4D4BA&usg=AFQjCNFaObAi5_3NNOdt8e1DiRoiHzg9bg&bvm=bv .47008514,d.dmQ ). To quote:

"Hydrogen peroxide with lowest recorded redox
potential of - 0.68 V compared to that of Cu++ / Cu+, +
0.15 V15 acts as a strong reducing agent particularly in
presence of hydroxide ions [13], [18] to donate electrons to
copper (II) forming copper (I) oxide,

H2O2 + 2 OH- → 2 H2O + O2 + 2 e- (1)
2 Cu++ + 2 e- + H2O2 → Cu2 O + H2O (2)

Reddish-yellow cuprous oxide is rendered colorless in
presence of sufficient ammonia to form
diamminecopper (I) [15],

Cu2 O + 2 NH4OH → 2 [Cu (NH3)2] OH + H2O (3)

[ not balanced, corrected per ajkoer:
Cu2O + 4 NH3 + H2O → 2 [Cu(NH3)2]OH (3)]

Diamminecopper (I), generated from reduction of
copper (II) or added exogenously facilitates oxidation
of ammonia, a reducing agent [14], by hydrogen
peroxide,

...[Catalyst].....Cu (NH3)2]OH.........................
NH3 + 3 H2O2 -----------------> HNO2 + 5 H2O (4)

[ not balanced, corrected by ajkoer:
NH3 + 3 H2O2 -----------------> HNO2 + 4 H2O (4)]

Further studies are required to elucidate the actual role
of diamminecopper (I) in the reaction; whether it is
converted to tetramminecopper (II), or undergoes a
reversible changes during the process."

With additional ammonia, the reaction with nitrous acid proceeds as follows:

HNO2 + NH3•H2O --> NH4NO2 + H2O

Interesting observations by the author's non-electrochemical experiment includes "The reaction is mediated by copper (II) as it fails to occur in absence of copper", and that the best order of addition of reactants is Cu then aqueous NH3 and finally H2O2. The author also notes the need for excess ammonia, to quote: "as it is needed to maintain: (i) solubility of copper; (ii) optimal alkalinity for expression of reducing potential of hydrogen peroxide; (iii) adequate concentration of free ammonia; and (iv) conversion of nitrous acid to ammonium nitrite."


Apparent weaknesses of this synthesis is that the key reaction of nitrite formation is a side reaction and, depending on reactant concentrations and time to addition of the H2SO4/FeSO4, yield could vary significantly. In addition, any NH4NO2 created is prone with time (possibly hours) to decomposition to N2, which could occur slowly, or possibly very rapidly including explosively especially on acidification of a concentrated solution (see Wikipedia on NH4NO2 at http://en.wikipedia.org/wiki/NH4NO2 and other MSDSs as well). The latter should restrict this preparation to small (or dilute) quantities only of HNO3.

Strengths of the synthesis include no need for a starting nitrate or nitrite salt or sophisticated equipment. The reactants are relatively available and inexpensive, and the reaction speed, depending on starting concentrations, is relatively rapid. There is also the possibility for alterations of the procedure to increase safety (like focusing on nitrate, and not nitrite, formation) and yield (like monitoring pH, varying and monitoring the temperature, concentration and reaction times including, for example, when to add Acetic acid/NaI, or H2SO4/FeSO4, for NO/NO2 creation). Employing electrolysis may also be able to increase yield and safety by increasing the formation of nitrates from nitirites by avoiding the addition of strong bases like NaOH. Source, here is an old paper on the electrolysis of aqueous ammonia in the presence of NaOH and, separately, Cu(OH)2 from 1905 (see page 242 at http://books.google.com/books?pg=PA242&lpg=PA242&dq=... from Journal Chemical Society, London, Volume 88, Part 2), to quote:

"Electrolytic Oxidation of Ammonia to Nitrites. Erich Muller and Fritz Spitzer (Ber., 1905, 38, 778—782. Compare Traube and Biltz, Abstr., 1904, ii, 727).—In the presence of a small amount of sodium hydroxide, ammonia may be oxidised electrolytically to nitrite even in the absence of copper compounds.

In the presence of copper hydroxide and sufficient alkali, the oxidation of ammonia to nitrite does not cease suddenly when the nitrite concentration has reached a certain value, but appears to proceed quite independently of the nitrite concentration. In these experiments, the oxidation was allowed to proceed for a comparatively short time only, so that the amount of alkali present was not greatly reduced. The formation of nitrite is intimately connected with the amount of alkali present, and when no sodium 'hydroxide is present, but only ammonia, nitrite, and copper hydroxide, it is found that the nitrite is transformed into nitrate more rapidly than the ammonia into nitrite, and thus the concentration of the nitrite tends to decrease.

Nitrogen is also formed during the oxidation. J. J. S."

Bottomline, while this synthesis may be doable as is, it more likely requires some work to address safety and yield considerations.


[Edited on 31-5-2013 by AJKOER]
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[*] posted on 31-5-2013 at 08:26


AJ, holy cow man the OP asked for minimum expense. With your procedure he is going to have to buy quite a few reagents.. There may be some place for these off the wall synths but I am beginning to get frustrated because they are in just about every thread and most the time not even what the OP wants.

I am sure you take great pride in your untested procedures and take the time to write them happily but It doesn't seem like anybody ever agrees with what you say or even tests them out. I am not trying to be mean but maybe you should invest your time performing these experiments you design instead of preaching them.

I 99.98% guarantee that a nitrate salt and H2SO4 is the cheapest method other than a bulk purchase of HNO3.




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[*] posted on 31-5-2013 at 09:05


Chemcam:

Thanks, I have revised my opening sentence as I do not make clear that my suggested synthesis chiefly addresses the issue of what to do with ones H2SO4 absent a nitrate or nitrite.

I also agree with your assessment that a nitrate salt and H2SO4 is the cheapest method assuming that one has the necessary equipment for the distillation. Those without such sophistications, adding Oxalic acid to NaNO3, IMHO, works much more easily and safely for a not too concentrated solutions (my recommendation based on working with H2C2O4 on acid salts to form various acids) for albeit a slightly greater expense.

Now, the availability and purity of the nitrate is, of course, still a possible issue for some.

As I completely disclosed with my thread extract, when I original peformed the first half of the reaction, the intent was not the formation of a nitrite or nitrate, but just a cupric salt. There is also a prior thread on Sciencemadness noting the formation of NH4NO2 (per the oxidation of aqueous ammonia) and testing it with H2SO4/FeSO4 that apparently forms some NO2 (yield is still an open question, please see http://www.sciencemadness.org/talk/post.php?action=reply&... to quote:

Quote: Originally posted by Formatik  
...
Quote:
Under what conditions can ammonia be oxidized to NH4NO2 ?


The Ber. ref. from Hoppe-Seyler describes it. Namely, strong solutions of H2O2 with a few drops of NH4OH or solutions of ammonium carbonate (with or without NaOH or Na2CO3) can be let to stand 24 hours without any nitrite formation occurring. But upon longer standing, even with a small amount of hydroxide then nitrite forms. Nitrite also forms when a dilute solution of H2O2 is mixed with NH4OH and a little Na2CO3 and is evaporated over pure conc. H2SO4 with a bell jar.

H2O2 forms (even in very dilute solutions) nitrite very rapidly, if the H2O2 solution is mixed with a few drops of NH4OH and a little NaOH or Na2CO3, and this then boiled in a retort to a very small volume. They suggest this nitrite formation as a demonstration experiment because it is very quick to do, and then after acidification of the colorless liquid with H2SO4, the HNO2 can be nicely be proven to be present.


So your implication that this is some entirely untried theorical based reaction is not quite correct. What I am guilty of, I fully admit, is perhaps overdoing the theory of the chemistry, and I do so when scholars will be reluctant to do so for reactions where the mechanics are not completely understood (as was the case for the action of aqueous ammonia and air on copper where there was at least 3 theories on the reaction path as of 1962).

Those interesting in exploring the extended procedure to HNO3 because of availability issues of nitrates, or commenting per their experience on suggested modifications, or theoretically finding the reaction interesting, are all welcomed. Personally, I find the very possibility of making Nitric acid from copper, household ammonia, dilute hydrogen peroxide, acetic acid and sodium iodide (with the latter two reactant, the focus is on nitrite formation, so also add a little NaOH) as just cool, irrespective of cost or yield.


[Edited on 1-6-2013 by AJKOER]
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[*] posted on 31-5-2013 at 14:02


He could also make it in a Birkeland-Eyde reactor using just air, water, and electricity.



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[*] posted on 31-5-2013 at 16:36


AJ, I was in a bad mood earlier from something else but I have since taken a chill pill. All I was meaning is that the OP asked for a simple and cheap way to make nitric acid. What you posted seemed much too complex and inefficient for someone who asks how to make HNO3 for the first time.

I have always had the impression that using copper metal as a reagent for making nitric acid is expensive and inefficient. Maybe when all nitrates have been banned your method will shine. And yes, someone who is looking for novel ways to HNO3 without a nitrate salt could benefit from what you said I just didn't think a beginner would go through all the trouble of acquiring the needed chemicals when easier ways exist.

Are you sure OTC 3% H2O2 and OTC 7% NH4OH w/ Cu will work? How fast is gas evolved? I would test it myself but I literally have no copper metal right now.




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[*] posted on 31-5-2013 at 17:31


Quote: Originally posted by chemcam  
....
Are you sure OTC 3% H2O2 and OTC 7% NH4OH w/ Cu will work? How fast is gas evolved? I would test it myself but I literally have no copper metal right now.


I used local US currency pennies which are actually Copper plated with a Zinc core. I am not worried about the Zn as I have read that a catalyst for the oxidation of ammonia by H2O2 one can use Zn powder, NaOH, Na2CO3,...

I could describe the reaction between Cu and dilute ammonia and H2O2, but best believed with ones own eyes. The nice feature of employing Cu is the reaction forms a colorful azure solution.

Note, for those wishing to comply with local laws (as I do), removing and changing coins avoids any Zinc contamination, preserves the coins and thoroughly cleans them.

Please read all the my caution warnings. Ammonium nitrite is also actually quite toxic on ingestion.


[Edited on 1-6-2013 by AJKOER]
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[*] posted on 1-6-2013 at 05:43


Here is a flowchart showing all the reasonable way to make nitric acid I come up with. It may help people stop asking questions about "how to make nitric acid":





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[*] posted on 1-6-2013 at 06:44


Plante1999:

May I suggested you add the general combustion of nitrogen containing fuels, as these can (and unfortunately often) form NOx exhaust.

Obviously, no commerical value as a path to HNO3. However, if a home chemist already has a sunk cost in an old lawn mover engine, for example, adding a green fuel may be a source of NOx.
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[*] posted on 1-6-2013 at 08:13


I do not think that is something to add to the diagram. Burning some fuel might indeed produce some NOx, but there will be sooo much other crap in the exhaust gases as well, that it is not economic and practical at all to isolate NOx and/or HNO3 from this. If that indeed were a viable route to HNO3 then many people would have done so, because of the general availability of the fuels.



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[*] posted on 2-6-2013 at 08:07


Woelen:

While I agree with your comments, my statement was quite non-specific, as my understanding is that NOx can be formed from the combustion of nitrogen containing fuels under certain conditions. Already cited above include the oxidation of urea and ammonia in the presence of catalysts. However, for example, when NH3 is combusted in a high pressure environment (like in an automotive engine), some NOx exhaust is formed without a catalyst.

The reason is, for example:

4 NH3 (g) + 5 O2 (g) ---> 4 NO (g) + 3 H2O (g)

so 9 moles are turned into 7, whereas:

4 NH3 (g) + 3 O2 (g) --> 2 N2 (g) + 6 H2O (g)

or 7 moles forms 8 moles. So in a high pressure environment, the first reaction may be favored.

Now, my suggestion implies that there could exist other simple nitrogen, hydrogen and other element compounds that when combusted at high pressure could similarly form NOx gas. No claim as to commercial significance, as this is not the topic. The current thread relates to restricted access to HNO3, and not a competitive cost effective, or even particularly safe, alternative route to Nitric acid.
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[*] posted on 2-6-2013 at 08:42


Can you get sulfuric acid (98%) and potassium, sodium, or ammonium nitrate?

If so...

Get a large beaker, 600 ml or so and place a fluted shot glass in the center. The stem on the shot glass should be several centimeters allowing the inside of the glass to be raised off the bottom of the beaker. Find some nitric acid proof plastic. I forget which plastic I used but I am pretty sure it was a ziploc baggie. Mix sulfuric and nitrate in the base of the 600 ml beaker and invert a corner of the baggie over the shot glass. Secure the baggie with a rubber band around the beaker, leaving a small hole over the beaker spout to allow for pressure release. Put ice and water in the baggie.

Put the contraption in a 60 C water bath for several hours, and when you come back, the shot glass will have filled with 100% fuming nitric acid, and if you keep the temperature low enough, it is generally mostly free of nitric oxides.
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[*] posted on 2-6-2013 at 08:51


Quote: Originally posted by binaryclock  
You really want nitric acid? That is some of the nastiest stuff I've ever encountered.

....

Anyways, see nurdrage's video regarding nitric acid if you really still want to make it: http://www.youtube.com/watch?v=2yE7v4wkuZU





You have not encountered too much then ;)

Do not make nitric acid like nerdrage did. He used HCl which contaminated his NO2 with nitrosyl chloride giving an impure nitric acid.
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[*] posted on 5-6-2013 at 09:05


Quote: Originally posted by The_Davster  


You have not encountered too much then ;)


I have not :).. pretty new to all of this. But nitric acid is vile stuff when it's fuming :)





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