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YT, I have to disagree with you (and hence agree with budullewraagh) on the KHSO4 and even more on the H3PO4 that you can use all equivalents for making HNO3. KHSO4 and KH2PO4 (and certainly K2HPO4) are not that strong acids and these are solids. With these, it will be hard (not impossible) to make HNO3, because you need strong heating and that will result in decomposition of any HNO3 formed into NO2, NO and O2. If you add some water, then you might be able to distill off an azeotropic mix, with KHSO4, but not with K2HPO4. The latter already is basic and will consume acid, instead of release acid. H3PO4 already is a weak acid, KH2PO4 is a VERY weak acid, K2HPO4 is a weak base, and K3PO4 is quite a strong base.

 

well I have a 14 month old daughter now, and she can join me in saying the first 12 elements (in order) and will often chunter away to herself the same :)

and yes, she has all the chems she`ll ever possibly want AND books, I`ve been stocking up on each :))

LOL :):D You indeed are a good dad! Other's first learn 1, 2, 3 or something like dad, mum, <names of other family members>, or whatever, you learn the periodic table. I like this :) . I hope she will really be interested in science when she grows up, but when you show your enthusiasm then that certainly works in the right direction. My youngest daugther (now 8 years old) is quite interested in all things I do, she especially asks not only about the "what", but also about the "why", she wants to understand things. If she may experiment herself, then she most likes it, but coming back to the topic, making HNO3 is not the thing she is allowed to do (at least not yet ;) ).

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before I had sulphuric acid, I had NaHSO4 from a basic childrens chem set, that with NaNO3 (and a little water) made Nitric acid, I have no reason to beleive the same should not (or would not) apply to a K salt also.

 

the H3PO4 and AN has made nitric acid here too, I sit less than 8 foot away from a test tube of it :)

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before I had sulphuric acid, I had NaHSO4 from a basic childrens chem set, that with NaNO3 (and a little water) made Nitric acid, I have no reason to beleive the same should not (or would not) apply to a K salt also.

Ah.. you added water. In that case, I can agree, the solution can be regarded as containing dilute HNO3 and distilling this can give some (dilute, at best azeotropic) HNO3. This is what I mentioned in the previous post. Making concentrated (90+ %) HNO3 from NaHSO4 (or KHSO4) and NaNO3/KNO3 will be MUCH harder.

 

the H3PO4 and AN has made nitric acid here too, I sit less than 8 foot away from a test tube of it :)

Yes, H3PO4 and AN will make nitric acid, but the lower phosphoric acids (NaH2Po4 and Na2HPO4) will make nitric acid with greater difficultly.

 

I even expect that if you mix AN and Na2HPO4, that the Na2HPO4 even acts as a base and NH3 is released, while NaNO3 and NaH2PO4 are formed.

EDIT: Fixed quote/quote error

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No it isn't. Copper nitrate looses oxygen as well, you get a mix of NO2, O2 and water vapor. When you lead this through water, you'll get nitrous acid, nitric acid, and unreacted NO2. The latter is very nasty. At higher concentrations, your acid will become yellow, brown, or green, due to the NO2/HNO2/N2O3 impurities. Such nitric acid is crap and cannot be used for many more sensitive reactions, because the nitrous acid and nitrogen oxides give very strong side reactrions.

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Do all Nitrates turn in to NO2 and O2? If you use Ammonium Nitrate will you get Ammonia?

The transition metal nitrates all decompose to the metal oxide, NO2 and O2.

The earth alkali metal nitrates also decompose to the metal oxide, NO2 and O2, but partly also to metal nitrites.

Lanthanide nitrates apparently do not decompose, I have cerium (III) nitrate n-hydrate, and it can be heated until it melts and gives no more water, but not a trace of NO2.

Alkali metal nitrates decompose to the nitrite and O2.

 

Ammonium nitrate is a special case, it decomposes, giving mainly water and N2O, but there also are side reactions with formation of NO, N2. If it is heated to strongly, then it may explode.

 

Nitrate salts of organic cations (e.g. subsituted ammonia salts, protonated amines) usually are very dangerous on heating. They have the oxidizer and reductor in the same compound. They can decompose very violently or explode. Many of them indeed are used as explosive. I once made formamidine disulfide dinitrate, this is the salt of protonated formamidine disulfide, and this compound even can ignite at room temperature (which I noticed, while I stored some of this and in the middle of the night it started burning spontaneously :eek: and the fire alarm went on).

 

Nitrate esters also usually decompose violently on heating. Many of them are suitable as explosives.

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Lithium nitrate fully decomposes to the oxide quite easily, the other group ones tend to go through the nitrITE stage and then with greater heat will decompose further.

 

Woelen, I made Cerium nitrate also, soaked in cloth, and then lit, it Did burn and left a mesh that glows Very Brightly over a gas flame, very similar to the Thorium nitrate gas Mantles :)

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  • 6 months later...

lance got nitric acid. heat egual amounts of potasium nitrate and sodium bisulfate (sodium hydrogen sulfate) together to get NO2. then just transfer the NO2 to some water and viola, you got nitric acid. the NO2 dissolves very easily in the water.

some caveats: use glass stoppers, NO2 will dissolve rubber ( i kow from experience) and NO2 is a pullatant and 2 grams will kill you

 

have fun

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NO2 + H2O will not give you nitric acid. It will give you mostly nitrous acid with perhaps a teency, tiny bit of nitric acid. Just look at the balanced equation and you'll see your problem.

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Jdurg, I can only partly agree with you. Making HNO3 of not too high a concentration can be done with NO2. When NO2 dissolves in water, it disproportionates to nitrogen(III) and nitrogen(V). The nitrogen(III) in turn gives rise to formation of NO and NO2.

 

Have you ever tried shaking NO2 with water? I did, in a sealed vessel. The deep brown gas quickly dissolves and what remains is a colorless liquid and a colorless gas. As soon as the gas is allowed to make contact with air, thick brown fumes are produced again.

 

The reaction is:

 

2NO2 + 2H2O <---> HNO2 + HNO3

HNO3 <---> H(+) + NO3(-) (very strongly to the right, so HNO3 is taken away from the system)

2HNO2 <---> N2O3 + H2O

N2O3 <---> NO + NO2

 

NO escapes as colorless gas, it hardly dissolves in water.

 

So, the net reaction is:

 

3NO2 + H2O --> 2H(+) + 2NO3(-) + NO

 

The NO can be reused, by allowing air to enter the mix. Then quick complete conversion to nitric acid occurs, because NO reacts easily with air to make NO2.

 

--------------------------------

 

Now the part, why I still do partly agree with Jdurg. The story, given above, only holds for low concentrations of HNO3. As soon as the concentration of HNO3 rises, then the NO2 remains dissolved. The ionization of HNO3 is not complete anymore. NO2 remains in solution, N2O3 also remains in solution, the liquid becomes brown/yellow, or even green, due to the blue color of N2O3. So, when one continues bubbling NO2 through water for a long time, then one finally ends up with acid of a few tens of percent (30%, 40% I do not know exactly), with lots of dissolved NO2 and N2O3, which is heavily giving off red vapor on contact with air. So, this method of making nitric acid only is suitable for making dilute to moderately concentrated acid.

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You need to add that after. If you add it before, the urea will destroy the NO2 already, before any acid is formed.

 

That method is nice for removing small amounts of NO2 (somewhat yellow HNO3), but a dark brown/red or green/yellow heavily fuming liquid is not something you want to clean with carbamide (urea). The urea usually contains other impurities and with the urea you also add impurities. If only small amounts of NO2 are in the acid, then that is not that bad, but with large amounts of NO2 you need large amounts of urea.

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you said "The NO can be reused, by allowing air to enter the mix. Then quick complete conversion to nitric acid occurs, because NO reacts easily with air to make NO2."

 

so would using H2O2 soln assist this rxn at all?

 

considering that Carbamide is a reductor, and H2O2 an oxidizer.

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I expect that H2O2 oxidizes any NO2 to HNO3 directly, so that would facilitate the reaction, but I'm not sure what happens at higher concentrations of HNO3. Sometimes, H2O2 also can work as reductor and then it could reduce HNO3 back to NO2, itself being oxidized to O2.

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  • 4 weeks later...

Read a few posts above your post, and then you can read that it is possible to make HNO3 from NO2 + H2O, but there are some issues with that. I explained in more depth in post #162, so please read that (again??).

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  • 1 year later...

I would like to decompose copper (II) nitrate and bubble that through water. I have learned that nitric can only be distilled to a concentration of 68%. To break this azeotrope, I had the idea of putting the solution into a Hoffman Apparatus and decompose the water. Would this work?

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  • 6 months later...

I am very sorry to be digging up such an old thread but some question needed to be answered.

 

I have tried making Nitric Acid by the method of KNO3 + H2SO4 = KHSO4 + HNO3

 

Is this procedure correct?

 

1) Boil the H2SO4 + KNO3 solution to HNO3's boiling point 83 °C

2) Bubble this gas into Distilled H2O

3) Keep the condenser unit running and keep the distillate under ice bath.

 

Also, when doing this distillation, instead of getting yellowish orange fumes, I get a very heavy orange to brown fumes that mask inside my flask. Here's a picture of it bellow.

 

Thanks,

Tim

NitricAcidFumes.jpg

Edited by hydraliskdragon
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You have some nitric acid, but there are nitrogen oxides present as well, "red-fuming" nitric acid.

 

Your ultimate goal should not be to buble a gas through water, but instead completely condense the gas into a liquid with as little water as possible. How is your condenser? I would recommend a graham condenser and some jointed glassware if you're really serious about this, but the key really is to keep your condenser really cold so you don't have so much gas. It is smart however to have some water in your recieving flask to dissolve the begginings of the acid.

 

I just found a good video, if you have the ability to distill under a vaccum, that would be a good idea, or a pretty cheap vaccum can be bought on consolidated-chemical.com.

 

Edited by frosch45
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Well, I figured it was nitrogen oxide a long time ago but was unsure why so much was produced at this stage.

 

Also, couldn't I simply use a liebig condenser?

 

Would a bath of ice and acetone in a Cold Finger condenser be suffice for the condensation?

Edited by hydraliskdragon
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Liebigs you can use, its just that they are less effective. make sure your liquid is very very very cold.

 

did you mean dry ice in acetone? that would be the very best that you are able to easily obtain

 

Also, in your picture, it looks like you are using a rubber stopper? I hope this isn't the case (or otherwise this is not inside your house) because if that is rubber, you better get outside of your house now and make sure all the windows are open

"Vapors are a strong irritant to the pulmonary tract. Initial symptoms of inhalation may be moderate and include

irritation of the eyes and throat, tightness of the chest, headache, nausea and gradual loss of strength. Severe

symptoms may be delayed (possibly for 5 to 7 hours) and include cyanosis, increased difficulty in breathing,

irregular respiration, lassitude and possible eventual death due to pulmonary edema in untreated cases."

 

http://www.vngas.com/pdf/g61.pdf

 

that stuff is really dangerous, and nothing less than well sealed glass or teflon is going to do much good

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