Sciencemadness Discussion Board

Nitric Acid Synthesis

 Pages:  1  2    4  

DJF90 - 7-9-2009 at 16:31

That would insult about 80% of the members of this board. Think before you speak. And you fall far short of a chemist.

entropy51 - 7-9-2009 at 16:32

Quote: Originally posted by Picric-A  


Im a chemist- i prefer to buy Reagent grade 98% sulphuric acid or 69% Nitric acid if i want it, not crappy 'drain opener' like you probably do!


If you can buy nitric (or sulfuric) then why on earth are you trying to decompose fertilizer (is that reagent grade too)?

And of course with your official police sanctioned research lab license you are allowed to possess "any chemical you desire".

If you're a chemist, why are you buying obsolete Rotovaps that don't even have half the parts?

If you're a chemist, why didn't you know the answer to this?


Quote: Originally posted by Picric-A  
can ammoniium nitrate be used to make nitric acid via H2SO4 + nitrate salt method?
Picric a

If you're a chemist I wonder why most of your posts read like they were written by a 14 year old kewl?

[Edited on 8-9-2009 by entropy51]

[Edited on 8-9-2009 by entropy51]

[Edited on 8-9-2009 by entropy51]

[Edited on 8-9-2009 by entropy51]

panziandi - 7-9-2009 at 16:37

Wait a minute there!

Quote: Originally posted by Picric-A  
If you have easy acess to conc sulphuric this reaction works extremly well.
Simply add pices of copper to a flask followed by conc sulphuric acid and simlpy heat to produce a steady stream of SO2, quite dry. This method works well when you have plenty of drain opener. Of cours it is pointless when using it to make sulphuric acid via contact process... :P


This was posted by you Picric-A on 30-8-2008 at 10:33 AM, http://www.sciencemadness.org/talk/viewthread.php?tid=11101#... now then you imply using cheap drain unblocker over lab grade due to cost, that is EXACTLY what Entropy is suggesting by using a cheaper "OTC" grade of sulphuric acid to generate your own cheap nitric acid.

DJF90 - 7-9-2009 at 16:40

There is nothing wrong with drain cleaner quality sulfuric acid. If one has the determination and patience, and suitable apparatus it can be cleaned to an equal specification as 98% reagent grade, and even higher if necessary. For many things, it will suffice where concentrated sulfuric acid is needed.

entropy51 - 7-9-2009 at 16:42

A real chemist would know that some things don't require reagent grade chemicals, wouldn't he?

I save my ACS reagent grade chemicals for the purposes that require them.

And if one knows what he's buying, not all drain cleaners are crappy. Some are 96% and "technical" grade.

[Edited on 8-9-2009 by entropy51]

DJF90 - 7-9-2009 at 16:43

A real chemist would know alot of things...

kclo4 - 7-9-2009 at 16:47

The continuing conflict between several members and Picric A and his derivatives is degrading the quality of this forum. I think all who participate should at try to hold back on the comments a bit, though it is damn hard I'm sure.

Picric A, want to help drop your conflicts between members, by not replying to what you might believe are personal attacks? I can see how your post above was provoked, but why do you choose to keep it going? Please stop for the sake of the forum.

As for the others, it would also be a good thing to stop this BS. Seriously, how many threads have derailed and ended with discussing the quality of Picric A?



panziandi - 7-9-2009 at 16:53

My Picric acid is ACS reagent grade...

entropy51 - 7-9-2009 at 17:00

OK, here you used a "roaring bunsen flame"
Quote: Originally posted by Picric-A  
i used roaring bunsen flame which proceeded to completion (on 2g scale)


But here you say you heated it with nichrome wire:

Quote: Originally posted by Picric-A  

how does it work?? well you could not be lazy and read the thread, but i guess i could tell you. I heated a copper tube full of Ca(NO3))2 beads with Nichrome wire then passed the evolved gasses through a series of bubblers.


Not only are you a liar Picric/Saber/Labxyz, but you are not even a good liar. It's beyond me why you haven't been banned for crapping up our forum with your fairy tales.
Quote: Originally posted by Picric-A  
If you had read what i wrote carefully you would note that.

Note- The hint is when i said on 2g scale and 'full of nitrate beads'
If you did write carefully (and you certainly don't) we wouldn't need hints to decipher what you claim to have done.

[Edited on 8-9-2009 by entropy51]

Picric-A - 8-9-2009 at 01:13

Roaring bunsen on test tube scale and nichrome on copper tube scale. If you had read what i wrote carefully you would note that.

Note- The hint is when i said on 2g scale and 'full of nitrate beads'

[Edited on 8-9-2009 by Picric-A]

starman - 8-9-2009 at 21:04

Quote: Originally posted by kclo4  
The continuing conflict between several members and Picric A and his derivatives is degrading the quality of this forum. I think all who participate should at try to hold back on the comments a bit, though it is damn hard I'm sure.

Picric A, want to help drop your conflicts between members, by not replying to what you might believe are personal attacks? I can see how your post above was provoked, but why do you choose to keep it going? Please stop for the sake of the forum.

As for the others, it would also be a good thing to stop this BS. Seriously, how many threads have derailed and ended with discussing the quality of Picric A?



Ive got to agree.Thre kid is trying to impress with his obvious flights of fancy and fails to understand that it seriously impedes his ability to benefit from the knowledge of seasoned members and real chemists.
However sustained derisory commentary from those that completely overwhelm the young Walter Mitty both academically and rhetorically is painful to witness.

[Edited on 9-9-2009 by starman]

[Edited on 9-9-2009 by starman]

[Edited on 9-9-2009 by starman]

1281371269 - 9-9-2009 at 11:03

Eloquently put.

uchiacon - 9-9-2009 at 20:21

Somebody ban that f*ckin...

I think he's on bombshock rubbing his cock all over the threads.
"Lol, y wud u waste all dis munney makin nitric when you can by it soo easy?"

That aside, I'm 14, and I don't act like that faggot.
Why did we even bother with this stupid decomp of nitrates for HNO3? Its dirty, annoying, with low yields and general faggotry. I would think that the bigger questions aren't being answered on this thread, but it seems people has plenty of time to argue wit that noob.

Would you guys think that a magnetic stirrer would be worth making to prevent bumpin on reduced pressure nitric distillations? Will bumping even occur at like 20-30 torr?
I know that a stirrer will increase how fast the nitric will evaporate, but is that worth integrating it into your home made heating mantle?

Oh, and I'd assume that if I manage to hook up two aspirators in series then I could increase the vacuum? I saw this on the other threads, but does anyone know by how much? i.e. two aspirators in series both capable for 29.5torr

Cheers

bilcksneatff - 10-9-2009 at 04:33

Quote: Originally posted by panziandi  
OR saturated calcium nitrate solution and add sulphuric acid (fairly strong if not concentrated) that would ppt CaSO4 and leave with a fairly storng nitric acid solution which could be filtered and distilled.


I think CaSO4 is pretty soluble in concentrated nitric acid, so this would only work form making dilute acid.

entropy51 - 10-9-2009 at 11:16

Quote: Originally posted by uchiacon  

Oh, and I'd assume that if I manage to hook up two aspirators in series then I could increase the vacuum? I saw this on the other threads, but does anyone know by how much? i.e. two aspirators in series both capable for 29.5torr


That would not be a correct assumption. The ultimate vacuum is set by the vapor pressure of the water. You should reach that with a good aspirator, and hooking two in series will not give a better vacuum than one by itself.

I would never vacuum distill without stirring on account of the likely bumping. Sometimes they bump even WITH stirring.

You sure your aspirator is only pulling 29.5 torr? Right now my aspirator is pulling 15 torr (read off a Hg manometer) at a water temperature of 24 C (which is not very cold at all.) Did you buy the aspirator new? Plastic (and to a lesser extent) metal aspirators do wear out and don't pull as low. Glass aspirators last forever (unless you break them.)

[Edited on 10-9-2009 by entropy51]

[Edited on 10-9-2009 by entropy51]

uchiacon - 10-9-2009 at 21:02

Well I just looked at the aspirator and it said 29.5 torr max... couldn't find many glass aspirators either. But wouldn't a metal aspirator be just as good as a glass one? They're both pretty darn hard.

This is one of the only glass aspirators I found, and it looks like it has that old style aspirator make. are these ones as affective as the normal ones?
www.onlinesciencemall.com/Shop/Control/Product/fp/vpid/17874...

And one from sigma aldrich, somehow at a crazy 150USD.
www.sigmaaldrich.com/catalog/ProductDetail.do?D7=0&N5=SE...

If you has any other sources for a normal looking glass aspirator do tell.

Cheers

entropy51 - 11-9-2009 at 05:54

I think the metal aspirators may not hold up in the long run if you distill a lot of HNO3 with them. HCl fumes probably won't do them any good either. Occasional use with acids may be OK, though.

I believe my glass aspirators were made by Kimble, but they look very much like the one in the Aldrich link above. Yes, they look old fashioned, but my glass aspirators wotk better than the metal and plastic aspirators that I've used. The vacuum is as good or better and it seems to me that the glass aspirators use less water, although I haven't measured the water consumption.

But if the one you have works for you then what more can you ask?

Magpie - 11-9-2009 at 07:54

Introduction of a tiny amount of air below the surface of the pot liquid will prevent bumping. I do this with a piece of glass tubing drawn to a pipette end. On the other end of the tube is a piece of rubber tubing with a screwed pinch clamp for adjusting the air flow rate. The tube is held in place with a thermometer adapter in a Claisen adapter. This is the method recommended in my organic chemistry lab manual. It works well.

If requested I can post a picture of this "ebulliator."

[Edited on 11-9-2009 by Magpie]

entropy51 - 11-9-2009 at 09:06

Magpie, I'll bet you and I are the only ones hereabouts who've used an ebulliator.

Gatterman would be proud of you!

I'm old school but I gave them up when I got a magnetic stirrer. They do work - unless solid forms in the pot and then I didn't find them as helpful.

dann2 - 11-9-2009 at 10:22


Will a few boiling chips not help here? Some pieces of red flowerpots sounds good.
Dann2

entropy51 - 11-9-2009 at 10:34

Boiling chips seem to lose their effectiveness under vacuum, although some people do use them. And bumping tends to be worse under vacuum. Boiling chips work fine for distillation at atmospheric pressure.

uchiacon - 12-9-2009 at 00:35

The metal aspirators are made out of brass with a laquer coating.
So not just plain metal.

Also tried out my 68% nitric today on some copper plating stripped off gouging rods, made tons of NO2. Lucky I did it outside, and then I reduced the copper nitrate to copper bits with zinc battery casings. Wish we could do that at school..

Can you give me a link for your kimble glass aspirator? Or any other links for a glass aspirator? And do the different designs make a difference?

Cheers

not_important - 12-9-2009 at 01:38

Quote:
Can you give me a link for your kimble glass aspirator? Or any other links for a glass aspirator?


http://208.72.236.210/html/pg-924000.html

http://www.onlinesciencemall.com/Shop/Control/Product/fp/vpi...


chloric1 - 12-9-2009 at 05:28

Quote: Originally posted by uchiacon  

Lucky I did it outside, and then I reduced the copper nitrate to copper bits with zinc battery casings. Wish we could do that at school.
Cheers

Man you MUST be joking I took highschool chemistry in 1990 and copper in nitric acid was an essential lab to demonstrate metal chemistry. We put copper in nitric acid in a fume hood, went home and then returned the next day and add water if needed and added sodium hydroxide to make copper hydroxide gell(loved that):D and then heated the gell to form black Copper oxide. After decantation 6M HCl was added to form clear green solution and an aluminum wire was dropped in. Went home and came back next day to find clear solution and salmon colored copper deposits with a little debris from the aluminum.

I suppose you are now going to tell me they no longer do the sugar in sulfuric acid trick or the lead iodide gold trick.

bilcksneatff - 12-9-2009 at 11:50

Quote: Originally posted by chloric1  
Man you MUST be joking I took highschool chemistry in 1990 and copper in nitric acid was an essential lab to demonstrate metal chemistry. We put copper in nitric acid in a fume hood, went home and then returned the next day and add water if needed and added sodium hydroxide to make copper hydroxide gell(loved that):D and then heated the gell to form black Copper oxide. After decantation 6M HCl was added to form clear green solution and an aluminum wire was dropped in. Went home and came back next day to find clear solution and salmon colored copper deposits with a little debris from the aluminum.

I suppose you are now going to tell me they no longer do the sugar in sulfuric acid trick or the lead iodide gold trick.


I took HS chem last year. We did none of those things! I wish we would have! I didn't like the teacher though; most of the class was bookwork and we hardly ever did labs. The most "dangerous" thing we ever did was titration of HCl with NaOH (and we did that within the first two weeks of AP Bio this year!). No nitric or sulfuric at all; even the teacher didn't demonstrate anything with them. We did maybe one metathesis, and it wasn't nearly as cool as the lead(II) iodide precipitation. Didn't even cover redox in the class. The worst part was that my school didn't have AP Chem, so I couldn't take a decent chem class!

Even so, thats what a home lab is for!

[Edited on 12-9-2009 by bilcksneatff]

[Edited on 12-9-2009 by bilcksneatff]

[Edited on 12-9-2009 by bilcksneatff]

1281371269 - 12-9-2009 at 12:02

We did pretty dull stuff in GCSE chemistry, I think the most exciting was heating Mg in a crucible until it popped, that was back when we were about 13. We did H2SO4 / Sucrose in triple i.e. for a class of about three people. But supposedly A level stuff is a lot more exciting (I've had three classes so far).

chloric1 - 12-9-2009 at 14:48

UGHH!! I was considering going back to college and just getting a Chemical Engineering degree. Man, I do not know if I can survive nothing but bookwork!! But my ebay is doing well and I should have much better home chemical setup in the not so distant future!:D

uchiacon - 12-9-2009 at 16:44

Chemistry in 1990 is a while back; things change in 10 years let alone 20.

The glass aspirator that was linked is of a different design( I linked it in my earlier post lol) and I asked whether the different aspirator designs make a difference in vacuum pulled.
Does anyone know?

1281371269 - 12-9-2009 at 17:15

I wonder if someone could sort something out for me about this aspirator stuff:
I've heard two descriptions of the way it works. The one that seems to make sense to me is that the tube is connected with an airtight seal to the top of a full pot of water. The water is run out. The air has to support the water and its pressure is greatly reduced, creating a soft vacuum.

The other one is about water running out at a certain speed somehow creating a vacuum...

I've done some searching on google etc, and it hasn't cleared up the issue.

ammonium isocyanate - 12-9-2009 at 17:50

It's called the Venturi effect.

Basically, a fluid moving at high speed pulls fluid around it into it's stream, propelling it in the same direction (this isn't really a very good description, I can post a sketch if necessary that explains it better). So in an aspirator, the moving water pulls air along with it, out of the chamber and into the atmosphere. Thus a vacuum is produced.

The same effect is observed in combo-type snowguns. Nucleation streams consisting of tiny ice particles are sucked up by water shot out of nozzles at high pressure, causing the water dropletsto freeze around the ice particles (if it's cold enough).

On the subject of aspirators, does anybody know how good the aspirators from pelletlab are? The're pretty cheap, so the're are great deal if they work well.

bilcksneatff - 13-9-2009 at 04:11

Quote: Originally posted by chloric1  
UGHH!! I was considering going back to college and just getting a Chemical Engineering degree. Man, I do not know if I can survive nothing but bookwork!! But my ebay is doing well and I should have much better home chemical setup in the not so distant future!:D


That was high school chem (In fact I'm still in high school, this is my last year). College chem should be much better than that ;)

[Edited on 13-9-2009 by bilcksneatff]

uchiacon - 13-9-2009 at 20:12

Guys, so what about this aspirator design www.pelletlab.com/filtering_kit

As compared to a standard design
www.carolina.com/product/filter+pump+or+aspirator.do?keyword...

Any difference?

Cheers

[Edited on 04-07-09 by uchiacon]

kclo4 - 13-9-2009 at 20:22

Quote: Originally posted by uchiacon  
Guys, so what about this aspirator design www.pelletlab.com/filtering_kit

As compared to a standard design
www.carolina.com/product/filter+pump+or+aspirator.do?keyword...

Any difference?

Cheers

[Edited on 04-07-09 by uchiacon]


Yes... they aren't even the same items. lol

One is a hand pump, the other is an aspirator...

uchiacon - 14-9-2009 at 00:54

Scroll down the page and click on the weird looking glass thing.

entropy51 - 14-9-2009 at 13:53

I've never used a glass aspirator of that design, but for $6.95 I doubt you can go wrong. The metal aspirator from Carolina seems quite reasonably priced and they claim it's very efficient. If I were in the market for an aspirator I would probably go for either one of those. The glass one may be easier to connect to your water supply unless you happen to have a faucet threaded to mate with the metal aspirator, but connections can usually be improvised with parts from your local hardware store.

uchiacon - 14-9-2009 at 20:39

heres a really cheap one from international peeps. $20US inc shipping. I think I might go for the metal one though..

www.onlinesciencemall.com/Shop/Control/fp/sret/1813212221162...

And for NZers, heres a place in Auckland that sells metal aspirators for $54 and $78.

www.deltaed.co.nz/

[Edited on 04-07-09 by uchiacon]

3287 - 8-12-2009 at 10:06

Was any progress made with producing N2O5 from NO2 + O3? It sounds like a remarkable way to produce strong nitric acid for those of us without access to nitrates or sulfuric acid.

dann2 - 20-12-2009 at 16:50

Hello,

2c worth on stuff that will/won't corrode in the presence of H2SO4/HNO3 (H2SO4/Nitrates too??) for making your own retort, or equivalent.
From this link:
http://www.corrosion-doctors.org/Why-Study/Right-material.ht...

Not too sure how to interpret the diagram. They say 'the shaded zones'. They should simply say, the zones? (five of them).
Zone 2 is what would interest the Nitric acid maker (via Sulphuric + Nitrate).
Cast iron pot/cauldron welded (using cast Iron rods) to an old cast Iron down pipe from spouting (that' guttering for the US audience :D)

Just realized, it says nothing about temperture.

Dann2

CALLING ALL MALE BEE'S, GET STIRRING THOSE CAULDRONS........

image015.jpg - 58kB

[Edited on 21-12-2009 by dann2]

hissingnoise - 20-12-2009 at 18:01

Silicon steel pots are used to recycle sulphuric acid from spent nitration acid---unfortunately they occasionally fail, and catastrophically, with the result that boiling H2SO4 enters the furnace. . .
I think I'll stick with glass; it's worked perfectly well, so far!





[Edited on 21-12-2009 by hissingnoise]

hissingnoise - 21-12-2009 at 08:25

Quote: Originally posted by 3287  
Was any progress made with producing N2O5 from NO2 + O3? It sounds like a remarkable way to produce strong nitric acid for those of us without access to nitrates or sulfuric acid.


N2O5 is produced electrolytically too; it's fairly complex but is cheaper than N2O4 oxidation by O3!
Producing dry O3 in sufficient quantity is problematic and collecting enough N2O4 isn't easy either.
If one had big bucks it could be done but there might be other, er, distractions if you had money to burn.
Buying out LLNL fr'instance---or making porn flicks!
But it looks like us mere mortals are stuck with KNO3/H2SO4 for the foreseeable future.
If you can't get those, you've got a real problem. . .



ScienceSquirrel - 21-12-2009 at 09:07

You can buy 38% nitric acid by the litre as pH Down in the UK.
It is used in hydroponics and is good to use as dilute nitric acid.

ScienceSquirrel - 21-12-2009 at 09:12

Quote: Originally posted by 3287  
Was any progress made with producing N2O5 from NO2 + O3? It sounds like a remarkable way to produce strong nitric acid for those of us without access to nitrates or sulfuric acid.


You would have to be a bit mad to go down this route. Both ozone and nitrogen dioxide are toxic and corrosive.
N2O5 is almost unbelievably nasty. It will explode on its own, it is corrosive and it is a very powerful oxidiser that will explode on contact with some materials.
At one time it was used as a nitrating agent but it has now been replaced by the far safer and more stable NO2 BF4.

hissingnoise - 21-12-2009 at 10:04

I know N2O5 will decompose quickly at room temp., I could be wrong, but I don't think the decomp. is explosive.
It's fairly stable in DCM and HNO3 and on its own at and below -60*C.
38% HNO3 needs a lot of H2SO4 if you want strong acid, but it's a lot better than nothing.

http://www.jstor.org/pss/53948


ScienceSquirrel - 21-12-2009 at 10:12

N2O5 is pretty safe in solution. The pure material is nasty.
It is just like nitrogen trichloride, pretty safe in solution in carbon tetrachloride, but a nice little yellow pool of the neat liquid is a definite no no....

hissingnoise - 21-12-2009 at 10:58

N2O5---> 2NO2 + O ---not a detonation; NCl3, though *does* detonate, sometimes apparently spontaneously.
Pure N2O5, itself, is a colourless crystalline salt.

ScienceSquirrel - 21-12-2009 at 17:49

Quote: Originally posted by hissingnoise  
N2O5---> 2NO2 + O ---not a detonation; NCl3, though *does* detonate, sometimes apparently spontaneously.
Pure N2O5, itself, is a colourless crystalline salt.


You mean;

2N205 -> 4NO2 + O2

Two moles of solid to five moles of hot gas, that is a detonation in anyone's book.
N2O5 runs around as NO2 NO3, the salt that you describe, but it also enjoys life as NO2ONO2.
Give the latter form a rubber crumb or two to bite on and it will take your hand off!

[Edited on 22-12-2009 by ScienceSquirrel]

hissingnoise - 22-12-2009 at 06:11

ScienceSquirrel, N2O5 is an oxidiser like N2O4 and is no more explosive than that substance.

ScienceSquirrel - 22-12-2009 at 08:40

A quick web search will reveal that N2O5 is a lot less stable than N2O4.
N2O4 is hypergolic with some hydrazines and will form explosive mixtures with some things.
N2O5 is in a different league, it explodes on contact with quite a few organic materials and some inorganic ones.
I cannot find a primary reference to it exploding on its own but there are plenty of secondary ones eg http://en.allexperts.com/e/n/ni/nitrogen.htm.
I would also point out that the structures are totally different, N2O4 is a dimer while N2O5 is not.
There are plenty of gotchas in chemistry where extrapolating the properties of one compound to another lands you in the doo doo.
Oxalyl chloride is often used as a model compound for phosgene and some of the chemistry is genuinely homogous but if you knew about oxalyl chlorides toxicity you just would not predict how toxic phosgene was.
Oxygen and ozone are both oxidisers but the difference of degree is huge. Condensed oxygen will just ignore a bit of silicon grease, condensed ozone will explode!

hissingnoise - 22-12-2009 at 08:53

I still contend that, by itself, N2O5 is not an explosive. . .
So, find a primary reference, and I'll concede the point!

ScienceSquirrel - 22-12-2009 at 09:07

I think we will have to agree to disagree, I have searched for a primary reference and failed to find one.
If I was planning on making it I would have a very careful literature search though.
One eye good, two eyes better :)

hissingnoise - 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.).

iHME - 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]

User - 23-12-2009 at 11:04

How is this relevant ?

hissingnoise - 23-12-2009 at 11:09

'Tis the season to be, er, tight?

iHME - 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. ;)

User - 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.

hissingnoise - 23-12-2009 at 14:22

I hope Iceland will accommodate all of us, because we're all coming, iHME.

kilowatt - 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?

arsen - 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?

chief - 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]

arsen - 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]

chief - 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 ...
=================

arsen - 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 ...
=================

Ephoton - 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]

not_important - 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.


Contrabasso - 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!

not_important - 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...


densest - 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.

watson.fawkes - 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.

chief - 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]

hissingnoise - 27-4-2010 at 08:55

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


The WiZard is In - 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!

arsen - 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]

chief - 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 ...

hissingnoise - 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.


chief - 28-4-2010 at 04:59

14 EUR ? Hopefully not for just 1 liter ??
==> H2SO4 can be got ..., just not from any store ..., except maybe a "drug-store", where it cost's 5 $/liter ...

Only such stores nearly dont't exist any more, and not each will sell it ...

I myself could get the 30-liter-canister for maybe 20 bucks, but I don't need it ...
==> Anyhow: When somehow roasting the plaster (CaSO4) it should be quite easy to get a bigger amount of the acid, the setup could be let run by itself, with not much attention ...

This way there would be a cheap source for other acids as well: HCl is obvious, but the usual standard-fertilizers could be distilled with the H2SO4 by the bag ... : Would certainly give quite a bit of HNO3 or maybe H3PO4 or whatever ...; this way re-crystallization could be avoided, the more simple distillation would come into it's place ...

Even the tubes of some old TVs could be used as a flask for that, if gently enough heated ...
==> With good thermal insulation and maybe the heat-input via a hot-air-gun from below (so they don't break such a process would require a day for a 25kg-charge, and electricity for maybe 1 or 2 $ ...)

Besides: In Germany, and probably many other states, almost no real chemical can be had from stores any more ...; everything thinned with loads of water, mixed with a ton of other ingredients and packaged in colorful bottles ...
==> Almost only useless crap in the stores ..., and much overpriced ...

hissingnoise - 28-4-2010 at 06:01

~chief, I'd happily pay more - making H2SO4 seems such a total hassle.
BTW, how come you're using dollars in Germany?



gnitseretni - 28-4-2010 at 06:06

Would heating the bejeesus out of a metal container filled with KNO3 over a fire and then bubbling the gases through water be a reasonable efficient way to make 60% or better HNO3?

If I'm not mistaken, first the KNO3 would decompose into KNO2 and O2. Then the KNO2 would decompose into K2O and NO and O2. The NO and O2 would react and form NO2 which when bubbled through the water forms HNO3.

What do you think? Worth a try?

KNO3 is the only nitrate I can get in large quantities cheap enough to try this.

chief - 28-4-2010 at 07:10

Depends on the fertilizer; but if you have clean KNO3, then it's a standard-textbook-way to make HNO3 by distilling it with H2SO4 ...
==> Some excess of H2SO4 is added, I believe the double amount ...,
==> and pure HNO3 distills over ...

I did it a couple of times with NH4NO3, following an old patent from 1911/1913, worked well, only I had use for gallons of the HNO3 and then ordered it, for 65 EUR-ct/kg ... :D:
==> It's said to be easier with NH4NO3 than with NaNO3 or KNO3, because everything in the flask is liquid and therefore no problems with partial overheating etc. occur.

================

When I mention prices here I convert them to $, because I can't find the EUR-sign on the keyboard ...

================

And now: What do most fertilizers consist of ?
==> probably no Chloride
==> maybe some sulfate, which doesn't react with H2SO4
==> maybe some phosphate
==> and maybe some nitrate ...

Don't know if any H3PO4 would distill over ...
==> also certainly not the H2SO4 ...

So what comes through the cooler would be HNO3, maybe H3PO4 (can it be distilled ?) ...
==> Might be a clean way to separate the precious parts from the fertilizer ...

=====================

Whoever suggests paying in excess of 15 $ for H2SO4/1 liter
==> is not a amateur-chemist,
==> but a hobby-chemist ... :o :D

[Edited on 28-4-2010 by chief]

hissingnoise - 28-4-2010 at 07:25

A well-stocked garden centre might have fertiliser marked 12-0-43 which is essentially pure KNO3. . .
And if you have HNO3, why not boil it with sulphur to get H2SO4?

[edit] Actually these two acids are soooo dependant on each other, it's sickening!
The bitches of chemistry?


[Edited on 28-4-2010 by hissingnoise]

Jor - 28-4-2010 at 07:39

If you boil HNO3 with sulfur, absorb the nitrogen oxides in cold water or better cold hydrogen peroxide, to regenerate nitric acid.

chief - 28-4-2010 at 08:41

Quote: Originally posted by hissingnoise  
A well-stocked garden centre might have fertiliser marked 12-0-43 which is essentially pure KNO3. . .
And if you have HNO3, why not boil it with sulphur to get H2SO4?

[edit] Actually these two acids are soooo dependant on each other, it's sickening!
The bitches of chemistry?


[Edited on 28-4-2010 by hissingnoise]


Never heared of that one ... ; what else can HNO3 be boiled with to give some acid ?
==> Anyhow the sulfur would be even harder to obtain than any of the previously mentioned ingredients ...

===============

I now wonder how the phase-diagram of HNO3/S/H2SO4 would behave under temperature ...
==> maybe 50%-HNO3 could be boiled with S to give H2SO4 with lower water-content, thereby generating higher-grade nitrating-acid by just boiling HNO3 with sulfur ??

Teh 69%-azetrope of HNO3/H2O surely is different in the presence of H2SO4, so that either in the boiling- or in the receiving- flask would be some higher-concentrated HNO3 possible ... ?

================

It also might cause thoughts of the reactins during the explsion of black-powder: The S-vapor might partially react with the NO2 of the nitrates to some H2SO4, which might play some role ... ... ?

[Edited on 28-4-2010 by chief]

hissingnoise - 28-4-2010 at 09:02

Quote:


It also might cause thoughts of the reactins during the explsion of black-powder: The S-vapor might partially react with the NO2 of the nitrates to some H2SO4, which might play some role?

The basis of the lead-chamber process, sans the charcoal. . .


gnitseretni - 28-4-2010 at 09:19

Quote: Originally posted by Jor  
If you boil HNO3 with sulfur, absorb the nitrogen oxides in cold water or better cold hydrogen peroxide, to regenerate nitric acid.


Instead of boiling HNO3 with sulfur, would heating the crap out of KNO3 work? That would produce oxides of nitrogen as well right?

hissingnoise - 28-4-2010 at 10:05

Thermal decomposition of KNO3 gives K2O, nitrogen and oxygen. . .
No NOx is formed!


gnitseretni - 28-4-2010 at 10:29

Well that sucks :(

hissingnoise - 28-4-2010 at 10:37

I know the nitrates of Ca and Cu do produce NO2 on strong heating. . .



gnitseretni - 28-4-2010 at 10:47

But I can't get these nitrates. (well not cheaply anyway)

I don't suppose one can make Ca or Cu nitrate from K nitrate?

chief - 28-4-2010 at 13:58

The nitrates of 2-valent Metals produce NOx upon thermal de-composition, those of 1-valent metals don't ...

arsen - 28-4-2010 at 19:38

Heating NH4NO3 has to be done very carefully, as the salt is prone to unstable decomposition, esp. if contaminated; thus synthesis via NH4NO3 is not recommended pathway.

Quote: Originally posted by chief  
Depends on the fertilizer; but if you have clean KNO3, then it's a standard-textbook-way to make HNO3 by distilling it with H2SO4 ...
==> Some excess of H2SO4 is added, I believe the double amount ...,
==> and pure HNO3 distills over ...

I did it a couple of times with NH4NO3, following an old patent from 1911/1913, worked well, only I had use for gallons of the HNO3 and then ordered it, for 65 EUR-ct/kg ... :D:
==> It's said to be easier with NH4NO3 than with NaNO3 or KNO3, because everything in the flask is liquid and therefore no problems with partial overheating etc. occur.

================

When I mention prices here I convert them to $, because I can't find the EUR-sign on the keyboard ...

================

And now: What do most fertilizers consist of ?
==> probably no Chloride
==> maybe some sulfate, which doesn't react with H2SO4
==> maybe some phosphate
==> and maybe some nitrate ...

Don't know if any H3PO4 would distill over ...
==> also certainly not the H2SO4 ...

So what comes through the cooler would be HNO3, maybe H3PO4 (can it be distilled ?) ...
==> Might be a clean way to separate the precious parts from the fertilizer ...

=====================

Whoever suggests paying in excess of 15 $ for H2SO4/1 liter
==> is not a amateur-chemist,
==> but a hobby-chemist ... :o :D

[Edited on 28-4-2010 by chief]

arsen - 28-4-2010 at 19:40

Solid amorphous sulfur is surprisingly has a quite inert aqueous chemistry, except in alkaline condition. I've never had any experience with this, but this pathway might not be as easy as it sounds.


Quote: Originally posted by hissingnoise  
A well-stocked garden centre might have fertiliser marked 12-0-43 which is essentially pure KNO3. . .
And if you have HNO3, why not boil it with sulphur to get H2SO4?

[edit] Actually these two acids are soooo dependant on each other, it's sickening!
The bitches of chemistry?


[Edited on 28-4-2010 by hissingnoise]

arsen - 28-4-2010 at 19:43

The main sulfur reaction during the combustion of blackpowder is an oxidation by KNO3 (and air if it's in open space) to form mainly SO2.
Quote: Originally posted by chief  
Quote: Originally posted by hissingnoise  
A well-stocked garden centre might have fertiliser marked 12-0-43 which is essentially pure KNO3. . .
And if you have HNO3, why not boil it with sulphur to get H2SO4?

[edit] Actually these two acids are soooo dependant on each other, it's sickening!
The bitches of chemistry?


[Edited on 28-4-2010 by hissingnoise]


Never heared of that one ... ; what else can HNO3 be boiled with to give some acid ?
==> Anyhow the sulfur would be even harder to obtain than any of the previously mentioned ingredients ...

===============

I now wonder how the phase-diagram of HNO3/S/H2SO4 would behave under temperature ...
==> maybe 50%-HNO3 could be boiled with S to give H2SO4 with lower water-content, thereby generating higher-grade nitrating-acid by just boiling HNO3 with sulfur ??

Teh 69%-azetrope of HNO3/H2O surely is different in the presence of H2SO4, so that either in the boiling- or in the receiving- flask would be some higher-concentrated HNO3 possible ... ?

================

It also might cause thoughts of the reactins during the explsion of black-powder: The S-vapor might partially react with the NO2 of the nitrates to some H2SO4, which might play some role ... ... ?

[Edited on 28-4-2010 by chief]

arsen - 28-4-2010 at 19:48

It is SAID, that most of transition metal nitrates liberate NOx upon heating, as well as alkaline earth metal but not alkali metal (except LiNO3, which is an anomaly).

Quote: Originally posted by chief  
The nitrates of 2-valent Metals produce NOx upon thermal de-composition, those of 1-valent metals don't ...

gnitseretni - 29-4-2010 at 06:28

I managed to find calcium nitrate. It's expensive though.. $45 for 50#. Geez!

How much HNO3 could I make from that? If I were to buy it, here's what I would probably try...

I would place the nitrate in a metal container. Then I'd make a hole in the lid for the Al tubing. The tubing would lead the gases into a 5 Gallon PP bucket filled with 2 gallons of cold water. I would put a lid on this bucket with a tiny hole in it so the NO2 doesn't readily escape but doesn't build up too much pressure either. And then start heating!

That's what I would do, unless someone has a better idea? But either way I'd have to think about it a little more cuz for $45... I dunno.. that's a little pricey!!

not_important - 29-4-2010 at 06:59

If you can get sulfuric acid, mix fairly dilute H2SO4 and a solution of calcium nitrate, separate the precipitate of hydrated calcium sulfate, concentrate the solution somewhat, then distill. When the original dilute HNO3 is 10% to 15% in strength it's not too bad for filtering or decanting, in terms of the bulk of the CaSO4 - which will be the hydrate CaSO4 2H2O Wash the ppt, use the wash liquor to make the next batch of calcium nitrate solution; use the CaSO4 to make desiccant by heating it to 150-180 C.

Going through the decomposition to NOx and forming HNO3 from that is a lot of work and somewhat wasteful as the nitrate has other decomposition paths besides giving NO2.

BTW, that calcium nitrate may well be a hydrate, so there's both less 'NO3' than you'd think, and heating it will result in the formation of a solution or mush of the nitrate and steam coming off before the nitrate decomposes.


Black powder - Potassium nitrate/sulphur

The WiZard is In - 29-4-2010 at 07:19

Quote: Originally posted by arsen  

It also might cause thoughts of the reactins during the explsion of black-powder: The S-vapor might partially react with the NO2 of the nitrates to some H2SO4, which might play some role ... ... ?


Extracted from :—

THE INITIATION, BURNING AND THERMAL DECOMPOSITION OF
GUNPOWDER
By J. D. BLACKWOOD AND F. P. BOWDEN, F.R.S.
Research Laboratory for the Physics and Chemistry of Surfaces,
Department of Physical Chemistry, University of Cambridge
Proceedings of the Royal Society (London)
Vol. 213. A. (8 July 1952) 285 1 19

(Received 3 November 1951-Revised 7 January 1952)

[Plates 3 to 6] {I do not have these plates. /djh/}

PART III. THERMAL DECOMPOSITION

It is well known that mixtures of potassium nitrate, charcoal and sulphur will react
exothermically with the production of a large volume of gas. These reactions
have formed the basis of extensive studies; one of the earliest and most
thorough is that of Noble & Abel. These workers examined the products but did
not formulate a reaction to explain the steps in the process. The more recent
contributions have been referred to earlier in the paper.

It appears that the presence of sulphur promotes reaction, and this has been
attributed by Hoffman (1929) to the formation of hydrogen sulphide at about
150oC from sulphur and organic matter present in the charcoal. This hydrogen
sulphide, he suggests, reacts exothermically with potassium nitrate above 280oC
to form potassium sulphate. It. is clear from the figures of Noble & Abel that the
amount of potassium sulphate formed is a function of the oxygen content of the
charcoal and not of the hydrogen content which only varies between 2 and 4 %.
It is also clear from experiment that the most reactive gunpowders are those
made from charcoal, the carbon content of which falls within quite narrow limits.
We shall see that reaction rate appears to be related to the organic materials
present in the charcoal as well as to the presence of sulphur, and the suggestion
put forward by Hoffman may require modification. Very little attention appears to
have been given to the preliminary reactions which appear to control the
behaviour of the gunpowder, and a study of these will be described in this
section.

The system potassium nitrate + sulphur

Experiments were carried out on mixtures of these materials. When they are
mixed together by grinding to a fine state, no reaction occurs. If they are heated,
gas evolution begins at a temperature of 250o C and is slow but continuous. As
the temperature is raised, the gas evolution is increased but is still of the same
form. Some typical results are shown in figure 15. The gases evolved during
heating are oxides of nitrogen, first nitric oxide and then nitrogen dioxide. If the
temperature is raised above the melting-point of potassium nitrate (334o C
sulphur dioxide can be detected quite readily although it cannot be detected at
lower temperatures. A small amount of potassium sulphate and potassium nitrite
is formed in the solid residue.

Attachment: Black Powder Blackwood & Bowden.pdf (327kB)
This file has been downloaded 887 times

gnitseretni - 29-4-2010 at 07:37

Ok, that's it! I will no longer waste time trying to find other ways to make HNO3. I'll just accept the fact that distilling is simply the best way to go!

Ah, I feel kinda relieved actually :P

chief - 29-4-2010 at 09:07

Quote: Originally posted by arsen  
Heating NH4NO3 has to be done very carefully, as the salt is prone to unstable decomposition, esp. if contaminated; thus synthesis via NH4NO3 is not recommended pathway.


Question would be: _How_ manageable is it ... ?
==> I didn't find the old patent yet ...

... but I did it: It was a setup involving a sand-bath .... in a remote room ... where noone was ... : I let it run for a time, then disconnected the power-source and looked at it later ...
If I remember right I even watched the temperature ... electronically ...

The NH4NO3 should clearly not be directly heated by a flame or whatever ...
==> Also the NH4NO3 I used was lab-grade ...
===================

I believe someone _here_ pointed me to the patent, in a similar thread ...
==> It was quite detailed about the temperature-control etc. ; still have it on my old laptop, will post it sometime ...

It was a german/austrian patent from 1911(Austria) and 1913 (Germany), shortly before the HNO3-synthesis from coal was invented ...
==> It also discussed the advantages over the then-standard NaNO3-way ... ...

====================

But one of my favourites, which I never tried, would be the use of either Ca(NO3)2 or Ba(NO3)2 with H2SO4:
==> Both form insoluble sulfates with H2SO4, so the Ba/Ca just ppt. as heavy=soluble sulfates, and the HNO3 can be just filtered or distilled ...
==> Question only would be how well it works with water-free H2SO4 to make dry HNO3: Will the HNO3 the come easily out of the sulfate-crystallizate ?

[Edited on 29-4-2010 by chief]

[Edited on 29-4-2010 by chief]

Mildronate - 29-4-2010 at 11:22

I learning chemistry in unversity and my university is a one man who working with glass he made for me retort.

Mildronate - 2-5-2010 at 09:25

Here you can see my retort, its made from 1liter flash and glass pipe. I use it only for nitric acid.
















 Pages:  1  2    4