Sciencemadness Discussion Board

Druken Aga Challenge (DAC) #3 - Closed (but open to discussion)

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Zombie - 5-3-2015 at 22:15

Good for you Delta!!!!!!!

I'm always gonna be a "billy"

I should'a thought that thru some...:(

deltaH - 5-3-2015 at 22:18

Thank you, thank you. Don't worry zombie, it's a fun reminder to all of us of the thumping thumper thread.

Zombie - 5-3-2015 at 22:28

It did sort of set the tone, and introduced my "personality"?.?.?

It's all good. Kudos!

blogfast25 - 6-3-2015 at 07:12

Quote: Originally posted by deltaH  
This is much of a muchness. Just pyrolyse the beans in a tin on the coals of a barbie and be done with it. Simple, no mess, no fuss. Don't forget to punch a small hole in the top of the tin for the volatiles to escape. You can then impregnate the charred beans with catalyst to your heart's content without anything going awry.



I'm a great fan of sacrificing quantity for better control. Something like the catalysed pyrolysis of dried beans carried out on a 100 g sample, with careful capture of any NO2 could teach us more than less controlled, larger trials.

Quote: Originally posted by Zombie  

I'm always gonna be a "billy"


Try and be aspirational? ;)




[Edited on 6-3-2015 by blogfast25]

Zombie - 6-3-2015 at 07:27

Quote: Originally posted by Zombie

I'm always gonna be a "billy"


Try and be aspirational? ;)


I'm going to buy a real friend today. :(

CuReUS - 12-3-2015 at 01:54

zombie's magnetron arc is amazing,even better than WGTR'S neodymium one
but I was wondering,why not use the arc to make things other than nitric acid,like HCN or ketene ?
acetylene +N2>(arc)>>2HCN

aga - 12-3-2015 at 12:14

Quote: Originally posted by CuReUS  
zombie's magnetron arc is amazing,even better than WGTR'S neodymium one

One small difference: WGTR's is Real.


[Edited on 12-3-2015 by aga]

Zombie - 12-3-2015 at 15:17

Quote: Originally posted by aga  
Quote: Originally posted by CuReUS  
zombie's magnetron arc is amazing,even better than WGTR'S neodymium one

One small difference: WGTR's is Real.


[Edited on 12-3-2015 by aga]
What on earth are you talking about?

I have burned more sh#t using magnatrons than any other plug in the wall device I can think of.

Set one of these inside the end of a pipe, and use resistance controlled SSR, as a controller.

You can set it to burn paper or just about melt the pipe.

Burning Bean Dust is a breaze.



[Edited on 3-12-2015 by Zombie]

aga - 12-3-2015 at 15:19

Cool !

Show us the photos and maybe win the challenge !

Edit:

Just collect the gas, and job done.

Hmm. Would the magnetron require/generate > 230V *1.41 ?

[Edited on 12-3-2015 by aga]

Zombie - 12-3-2015 at 17:23

Here's a couple of Utube vids that explain it.

There are two ways to control the output.
The first (better) is to interrupt the +5volt filament lead with the SSR. This limits the amount of electrons produced, and lowers the output power.

The second method is to interrupt the line voltage in (110V AC) with the SSR.
The issue there is the power output is not as easy to regulate. You go from nothing to near full power in an instant.
The video will explain why this happens.

I would build the entire magnatron, and at least 1 foot of 2 inch diameter pipe into a concrete insulator similar to a forced air furnace. It will get that hot if you allow it to. A cooling fan will go a long way to making this work long term.

https://www.youtube.com/watch?v=qT6EmMkKevY

https://www.youtube.com/watch?v=I2k2g00onL0



[Edited on 3-13-2015 by Zombie]


[Edited on 3-13-2015 by Zombie]

Untitled.jpg - 1.1MB

I just came across this article...

http://www.google.com/patents/WO2013141725A1?cl=en

[Edited on 3-13-2015 by Zombie]

[Edited on 3-13-2015 by Zombie]

CuReUS - 13-3-2015 at 01:18

Quote: Originally posted by Zombie  

I have burned more sh#t using magnatrons than any other plug in the wall device I can think of.
Burning Bean Dust is a breaze.

No,You don't understand.I meant putting the magnetron arc in WGTR's setup instead of the neodymium one,not using it to burn beans:mad:
I thought that idea crashed and burned long ago

j_sum1 - 13-3-2015 at 02:35

Hey. Burning beans hasn't crashed. They just haven't burned yet.
Maybe this weekend.

Zombie - 13-3-2015 at 07:41

Quote: Originally posted by CuReUS  
Quote: Originally posted by Zombie  

I have burned more sh#t using magnatrons than any other plug in the wall device I can think of.
Burning Bean Dust is a breaze.

No,You don't understand.I meant putting the magnetron arc in WGTR's setup instead of the neodymium one,not using it to burn beans:mad:
I thought that idea crashed and burned long ago



I don't know what WTGR's neodymium thing is. :(


Quote: Originally posted by j_sum1  
Hey. Burning beans hasn't crashed. They just haven't burned yet.
Maybe this weekend.



I'm sure you get it! :cool:

Adsorption and Desorption of Dilute NO2 on Silica Gel

WGTR - 15-3-2015 at 09:18

This is potentially important for anyone generating a dilute stream of NO2, regardless of how it is made. A dilute feed of (ideally) dry NO2 can be pumped through silica gel (indicating, in this case). Upon saturation, the beads can be heated up, causing them to desorb the concentrated NO2.

The following video is my "neodymium thingy" generating some NO2, as described in my earlier posts. NO is generated in the small arc chamber, which then oxidizes to NO2 in the bottle. The dilute NO2 is then adsorbed onto a column of indicating silica gel. The bubbler in the test tube functions just as a flow meter of sorts. All of the NO2 is adsorbed before it gets that far.

Attachment: NO2_Adsorption.avi (9.6MB)
This file has been downloaded 764 times

After placing the saturated beads into a sealed test tube, I heated the bottom with a heat gun to drive off some NO2:

Desorption.jpg - 577kB

I didn't heat the beads very much, because I didn't want to pop the stopper off the tube. I had to be careful doing this. In other words, the picture doesn't demonstrate very efficient desorption. Some moisture was adsorbed from the gas stream, shown by the condensation under the test tube stopper.

Zombie - 15-3-2015 at 09:30

That is just insanely creative.

Have you managed to create any Nitric acid from this yet?

WGTR - 15-3-2015 at 10:24

Small amounts. Right now I'm just trying to figure out what I'm doing. This particular setup is small, and might be able to make 1 ml of concentrated acid per day. The arc chamber itself is only using about 5W average power.

I'll probably add a moisture adsorption column right before the arc chamber, to keep the adsorbed NO2 dry...not that it matters very much. To scale this up, I'm planning on utilizing the tube furnace that I put together here:

https://www.sciencemadness.org/whisper/viewthread.php?tid=55...

When cool, silica gel in the tube could adsorb the dilute NO2. When heated, concentrated NO2 can be flushed out, regenerating the column. I didn't mention this previously, but it is much easier to make nitric acid from concentrated NO2 than it is from the diluted variety. NO oxidizes much faster to NO2 when it is concentrated. That's one reason I'm doing it this way. Another reason is that the creation of NO2 and the formation of nitric acid do not have to happen at the same time. Theoretically a bottle full of silica gel could store concentrated NO2 for a long time.

As an aside, only NO2 is adsorbed in the silica gel, not NO. This can be a useful way of purifying NO.

deltaH - 15-3-2015 at 11:15

That's very smart WGTR, genius. Well done indeed!

aga - 15-3-2015 at 11:48

Positive proof that chasing insane ideas can lead to insanely wonderful discoveries.

BRAVO WGTR !

deltaH - 15-3-2015 at 12:32

Can one not add the NO2 loaded silica straight to water or even better, dilute H2O2? If you had enough, then eluting through a short column could yield pretty strong acid perhaps.

WGTR - 15-3-2015 at 21:39

Quote: Originally posted by deltaH  
That's very smart WGTR, genius. Well done indeed!

Well, you know what they say...every family should have at least three children. That way if one of them turns out to be a genius, the other two can support him! :cool:

Thanks for the compliments all three of you. I do, however, stand on the shoulders of people much smarter than I am. I hope not to imply otherwise. The idea of NO2 adsorption onto silica gel is not a new one. Here is the patent that I borrowed the idea from:

Attachment: US2578674.pdf (1.8MB)
This file has been downloaded 742 times

Quote: Originally posted by deltaH  
Can one not add the NO2 loaded silica straight to water or even better, dilute H2O2? If you had enough, then eluting through a short column could yield pretty strong acid perhaps.

Those are interesting ideas. The first one would work if the gel is reusable, and doesn't disintegrate in the process. The second idea would work if one doesn't mind buying the peroxide from the grocery store.

Silica gel does a lot of interesting things to different substances. I don't claim to know much about it right now. Depending on whether it holds on better to water or nitric acid, it may be possible to concentrate it by passing nitric acid vapors over it. I seem to remember reading somewhere that it is very difficult to get all the nitric acid out of silica gel, even by rinsing it with water. I'll have to double-check that, though.

I'm starting school again tomorrow morning. That means that I'll be quite busy again, unfortunately. I cleaned out the fume hood a few hours ago to prepare for tomorrow's work. This project was carefully disassembled and put in the cabinet, where it will rest for hopefully not very long. Have a good week everybody.

CuReUS - 17-3-2015 at 00:59

Quote: Originally posted by deltaH  
Can one not add the NO2 loaded silica straight to water or even better, dilute H2O2?

Quote:
if one doesn't mind buying the peroxide from the grocery store

50% H2O2 can be made by electrolysing dilute H2SO4
WGTR,what about replacing that neodymium disk with a magnetron arc,or is it jus too dangerous ?
also,no one paid attention to this:(
Quote:
but I was wondering,why not use the arc to make things other than nitric acid,like HCN or ketene ?
acetylene +N2>(arc)>>2HCN


aga - 17-3-2015 at 14:34

Quote: Originally posted by CuReUS  
no one paid attention to this:(

Probably because the thread is about making HNO3.

Do some experiments with arcing and see what else it can be applied to.

WGTR - 18-3-2015 at 14:21

Quote: Originally posted by CuReUS  

WGTR,what about replacing that neodymium disk with a magnetron arc,or is it jus too dangerous ?
also,no one paid attention to this:(
Quote:
but I was wondering,why not use the arc to make things other than nitric acid,like HCN or ketene ?
acetylene +N2>(arc)>>2HCN



Since you ask, frankly, the magnetron idea frightens me, about as much as using a high-powered laser. One must be careful of any unintended reflections of energy in both cases.

I've tried to design the arc chamber such that it doesn't require a high voltage, high current power supply. I've separated the high voltage and high current functions into separate power supplies, such that there is no current behind the high voltage igniter, and there is relatively low voltage driving the high current arc. I actually got across the high voltage ignition coil. It was unpleasant, but still reasonably safe for a healthy individual.

This type of system is not intended to be something that a trained squirrel can operate. At the same time, I figure that some 12-year-olds somewhere may try building this. I'm trying to be careful not to give them a way to inadvertently kill themselves.

When I work with cyanides, I go to great lengths to avoid making HCN. Any time I have tried making cyanides, I did so without creating HCN itself. Ketene is something I've never tried making. It's hazardous, and I don't really need acetic anhydride for anything.

aga - 18-3-2015 at 15:45

Funny how people with Common Sense tend to live longer ...

Zombie - 18-3-2015 at 15:56

Languishing in misery ... :D

aga - 18-3-2015 at 16:00

Falsch.

People with a bit of Common Sense can enjoy (!) a bungee jump, but they check the Fall distance, extensibility and length of the bungee and their own weight before leaping off a bridge, hurtling towards the ground.

Zombie - 18-3-2015 at 16:52

Yet brilliant people will figure a way out of the dilemma before hitting the ground.

Who do you think had more fun? :o

Sorry... wrong thread.

WGTR - 18-3-2015 at 19:38

Just for fun, I tried something different. No pictures or videos this time; I was just doing a quick test to see what would happen with a certain procedure.

As derived from this reference, I tried electrolysing ammonia to see if it would oxidise at the anode. A nickel cathode and platinum anode were used for the test. The electrolysis cell was simply two 25mL beakers connected via a glass salt bridge. This configuration provided a separate anode and cathode compartment. 1g of NaOH was added to each beaker in 30mL of DI water. A few 10's of mg of CuSO4 was dissolved in 3.5mL of 30% ammonia, and the deep blue solution was added to the anode compartment. The electrodes were inserted, and electrolysis was maintained at 100mA for approximately 3 hours. Cell voltage was high, around 30V, due to the resistance caused by the salt bridge. No gasses were collected, but it appeared that the production of cathode gasses was very much greater than that of the anode gasses.

The cell was left in operation overnight, but the anolyte level dipped below the anode, breaking the connection. I estimate that the cell operated for a few hours. In the morning, the previously deep blue anolyte was clear, and the copper was precipitated as a black film at the bottom of the anode compartment. The nickel cathode was still bright and clean, with no obvious contamination. The pH was not measured in this test.

The anolyte was poured into an evaporation tray, and gently evaporated down on a hot plate. The resulting moist slush was added to a 20mL test tube, rinsed to the bottom with a small amount of DI water, and acidified with 98% H2SO4. The initial drop or two of acid caused a violent release of gasses, almost ejecting material from the tube. Further drops had little comparative effect. There was no noticeable color to the gasses.

Phosphorous-de-oxidized copper was filed into a fine powder, and a small amount was added to the test tube solution. Gentle heating was applied with a propane torch. The copper powder immediately began reacting, producing copious amounts of gas. Further heating was not necessary. The solution began turning blue, and significant amounts of brown gas filled the test tube. Strong bubbling continued for about 10 minutes.

It appears that nitrate was produced by the electrolytic oxidation of ammonia. A strong positive result for nitrates was given by the copper turning test. A nickel cathode was used due to its low over-potential for hydrogen production. A bright platinum anode was used in order to duplicate the original experiment. Also, it has a high over-potential for oxygen production, and is dimensionally stable under the cell conditions. It is possible that a different material, such as steel, could be used. This may be verified at a later time.

j_sum1 - 19-3-2015 at 02:16

Wow! Cool result.
I like this approach. Worth investigating further.

NWS - 12-6-2015 at 08:33

So I stumbled on this thread while looking for a way to make nitric acid after having used up all my OTC nitrates, and I have to say it's very impressive. Concerning the electrolysis, I actually happen to have some information that may or may not be helpful. The 8th edition of quantitative chemical analysis give the following reactions and E not values in volts.

HN3 +3H+ 2e- <--> N2 + NH4+ at 2.079 V (Avoid the reverse reaction at all costs!)
N2O(g)+ 2H+ + 2e- <-->N2(g) +h2O at 1.769 V
(laughing gas from nitrogen and water)
2NO(g) + 2H+ + 2e- <--> N2O+ H2O at 1.587 V
(reverse this and make NO from laughing gas?)
NO+ + e- <--> NO(g) at 1.46
(if someone finds a use for this then good for them)
HNO2 + H+ e- <--> NO(g) + H2O at 0.984 V
(same here)
NO3- + 4H+ + 3e- <--> NO(g)+ H2O at .955 V
(more or less useless because it'd probably be easier to oxidise NO to NO2)
NO3- + 3H+ +2e- <--> HNO2 + H2O at .940 V
(same here)
NO3- + 2H+ +e- <--> 1/2 N2O4 +H2O .798 V
(so avoid this because you'll just waste your precious nitrate)
N2+ 8H+ +6e- <--> 2NH4+ at 0.274 V
(a good way to get rid of ammonium if you need to)
N2+ 5H+ +4e- <--> N2H5+ at -0.214
(I don't like the looks of N2H5+)
3/2N2 + H+ + e- HN3 at -3.334
(avoid this one like the plague

and then for water
H2O + e- <--> 1/2H2 + OH- at -0.828 V
1/2O2 + 2H+ + 2e- <--> H2O at 1.229 V

It says nothing about ammonia/um to nitrate, but It might still be possible.

[Edited on 12-6-2015 by NWS]

blogfast25 - 12-6-2015 at 10:02

Quote: Originally posted by NWS  


It says nothing about ammonia/um to nitrate, but It might still be possible.


Of course it's possible, see e.g. nitrogen fixation:

https://en.wikipedia.org/wiki/Nitrogen_fixation

etoxiran - 12-6-2015 at 22:36

Hmmm, you said no Pt allowed. But remember that almost every car in EU/USA/Australia produced after 2007 must have DPF/FAP filter. Main principle of it's work was oxidizing soot to carbon dioxide, by NOx produced from exhaust gases over Pt catalyst.
High Pt load catalyst to be precise.
Cost of this part is about 50-250 euro, if you don't want disassemble your own vehicle, why not buy this part from ebay or local muffler shop?
It works (I don't have to use ideas like that because in my country I can buy HNO3 55% as cheap as 1 euro/dm3), but here you have ready to use apparatus (need only reconfigure a bit and add heating unit).
Virtually can be running using only heat & air. Adding to air inlet additional oxygen (welding store), will make more NOx.
http://www.ecovehicletuning.co.uk/images/dpf-overview.jpg

Before i forget - catalytic converter is unsuitable - because can consist rhodium particle which are reducing NOx to N2 & O2.




[Edited on 13-6-2015 by etoxiran]

NWS - 12-6-2015 at 23:04

http://digipac.ca/chemical/mtom/contents/chapter3/fritzhaber... Is this too good to be true? :o modify this so the copper is partially submerged in the ammonia, vaporizing more and preventing the copper from melting, Oxidize the NO to NO2, then store in silica or bubble though water?

gatosgr - 12-6-2015 at 23:13

What kind of apparratus do you need for this?

[Edited on 13-6-2015 by gatosgr]

NWS - 12-6-2015 at 23:32

I'd imagine one could redneck it pretty easy, suspend the copper wire over a 250 ml flask, use a stopper with a hole and then either some tubing to bubble your hot gas mix through water or perhaps run it though a condenser
edit- Household ammonia doesn't work, I'm going to try and prepare something a bit stronger by nutralizing some of the ammonia into ammonium chloride, drying, adding back to "regular" ammonia, adding a little NaOH and then distilling ammonia off the mix.

[Edited on 13-6-2015 by NWS]

gatosgr - 13-6-2015 at 01:25

You need 33% ammonia for this reaction. Nitric acid destroys rubber doesnt it?

[Edited on 13-6-2015 by gatosgr]

NWS - 13-6-2015 at 07:48

It does, however HDPE is fairly resistant to it, plus the extra ammonia fumes will probably offer some small protection.
Edit- where did you find that you needed 33%?

[Edited on 13-6-2015 by NWS]

blogfast25 - 13-6-2015 at 07:58

Quote: Originally posted by gatosgr  
Nitric acid destroys rubber doesnt it?



Too vague. Many, even affordable elastomers resist nitric acid very well. Ethylene propylene copolymers (EPR, also EPDM) for instance are comparable to (HD or LD)PE in their resistance to nitric acid. Peroxide-cured EPR is very resistant to it, even concentrated NA.

Hydrogenated nitrile rubber and hydrogenated SBS (SEBS) thermoplastic copolymer elastomers also exhibit good resistance to NA.

Quote: Originally posted by NWS  
It does, however HDPE is fairly resistant to it, [...]


Make that 'very', instead of 'fairly'. 75 w% HNO3 (and higher) is routinely packaged in HDPE containers. The one I buy is. There's nothing in NA that can oxidise the paraffinic structure of PE.

[Edited on 13-6-2015 by blogfast25]

gatosgr - 13-6-2015 at 08:10

From youtube search for the ostwald reaction.

[Edited on 13-6-2015 by gatosgr]

If anyone wants to participate you can make nitrogen by:

Nitrogen may be prepared from air by passing the air through a solution of caustic soda to remove carbon dioxide.

2 NaOH + CO2 ==> Na2CO3 + H2O

Then the gas stream is passed over heated copper turnings which removes the oxygen.

Cu + O2 ==> 2 CuO

Although I don't know what other steps you can take for making HNO3

[Edited on 13-6-2015 by gatosgr]

NWS - 13-6-2015 at 13:54

Wait, when did we go from ammonia to nitrogen? What is the nitrogen being used for?
also I stand corrected on the rubber and HDPE.

gatosgr - 13-6-2015 at 14:21

nitrogen and a kite for high voltage :D

http://link.springer.com/article/10.1007%2FBF00764903#page-1

to make ammonia since you can't buy it can you? or other nitrates e.t.c.

here dont ask again:

https://www.youtube.com/watch?v=CnhAnhK7U18

[Edited on 13-6-2015 by gatosgr]

NWS - 13-6-2015 at 15:27

So it looks like we're back to electricity/ distilling urine XD
and gotcha

j_sum1 - 13-6-2015 at 16:26

******************************************************************

Interesting that this thread is resurrected and gaining some new interest. It would be beneficial to newcomers to read it through. There is much discussion that has gone on in the preceding pages. A huge number of ideas have been thrashed out. Some still hold potential and some have been rejected. Here is a bit of a summary.

******************************************************************

1. As far as I know, the competition is still open. Better check with aga though.

2. Even for those who live in places where nitric acid and nitrates are difficult to procure, the simplest and cheapest method is still to purchase it somehow. Nothing posted in this threat has threatened that position.

3. I am not aware of anyone actively working on the project currently. I believe that there are a few shelved projects that have not been given up on.

4. To my knowledge there are two half completed projects that hold some potential:
(a) my own project based on the burning of plant proteins (lentils) to produce NOx which can then be absorbed by water to produce a very dilute acid that can be boiled down. I have redesigned my furnace and flue and am confident of getting something but expect yields to be very low and for it to be very time and energy consuming. My aim is scrap heap/ third world technology with zero dollars on consumables. But I have added copper + manganese catalyst.
(b) WGTR's variation of the Birkeland-Eyde process. He has built some impressive equipment but is yet to sort out the air flow. I expect yields again to be low. His is a lot more sophisticated than mine.

5. Other options
(a) the Ostwald process (ammonia + oxygen --> nitric acid) and variations thereof have to be rejected as outside the scope of the competition. The necessity for (i) high pressure, (ii) high temperature (iii) a platinum-rhodium catalyst and (iv) refeeding 95% of the product stream back through the process to lift yield above marginal values all combine to this being a highly technical approach. And that is before you consider the ammonia source material. (Of course if you can design a suitcase sized "black box" that you can feed ammonia in and get a trickle of nitric acid out, you will generate some interest and probably some sales.)
(b) The French method for producing nitrates and variations of that method are problematic due to the fact that these are (i) long-term projects (ii) requiring large volumes of material to be workable, (iii) taking up considerable space (ie, not an urban back yard), and (iv) having considerable smell. To my knowledge no one is currently working on this although If they had started back when the competition started they would well and truly have the jump on everyone else by now.
(c) Other novel ideas have been presented including using urea as feedstock, but no viable routes have been devised.

NWS - 13-6-2015 at 21:06

Well what I've been spit balling for a few days now is a modified otswald processes with copper as a catalyst, the only thing holding me back from actually trying is the need to make .880 ammonia, and yea, I've read pretty thoroughly looking for the magic bullet to make some nitric acid XD
edit- it may or may not end up that it simply makes ammonium nitrate, but I'd still be happy as can be if I got that far

[Edited on 14-6-2015 by NWS]

WGTR - 13-6-2015 at 21:15

I was thinking about this project yesterday. I haven't given up on it, but the realitites of life are such that I've had to focus on other things lately. The basic design was working, although there are still a few bugs to work out of it. Ironically, one thing that has slowed me down is that I've tried to make this a very simple design. I'm going to readjust my design a bit, and it may end up becoming more sophisticated than most people would like. I think this is necessary just in the interest of me actually finishing the project. We can always simplify things later.

I don't recall an announcement from aga stating that he had run out of beer, so I assume the competition is still going. I took myself out of the running for the prize, so I'm just doing this for the challenge.

deltaH - 19-6-2015 at 02:34

To expand on the electrolysis idea, earlier in the thread, I've mentioned a thesis that demonstrates that manganese dioxide is a functional low-cost ammonia oxidation catalyst. Fortuitously, MnO2 can also be used to make cheap amateur-friendly anodes (albeit not very durable) since it is an electrically conductive oxide. So...

BTW when run under basic conditions, you need added electrolyte (e.g. NaOH/KOH) because the ammonia does not ionise enough for good electrical conductivity, so sodium/potassium nitrate would be the product generated if successful :(

EDIT: Nurdrage's method for making MnO2 anodes:
https://www.youtube.com/watch?v=nvMVlhBmv7M

[Edited on 19-6-2015 by deltaH]

deltaH - 19-6-2015 at 03:09

I've just remembered an odd experiment I've done many years ago. I've dissolved some melamine (the chemical, not the plastic) in water (can't dissolve much, but slurry works), then slowly added spatula full amounts of calcium hypochlorite (as pool dry chlorine granules) with mixing and waiting for the hypochlorite to react between additions . The solution initially turned yellow, then on further addition turns a very beautiful intense cherry red, then if you continue even more additions, VIOLENTLY oxidises to generate copious amounts of a noxious NO2 smelling gas and turns clear again.

This experiment is maybe worth a revisit in light of this challenge as I think this may have been oxidising fully to make CO2 and NOx or nitrates. Feel free to do so, I cannot do it in my current circumstances. Perhaps melamine can be substituted for another readily available nitrogen source?

Since melamine is made from urea, urea might work as well. Although I don't suspect urea would form the beautifully coloured intermediates that melamine did. Be warned that this can potentially generate toxic and dangerous compounds besides the nitrogen oxides.

Urea granules + pool chlorine granules + a little water + induction period = CO2 + NO2 + crud + woosh?!

sounds a little too good to be true and dangerous :o

Ok, since this is a science forum, I had probably write that equation a little better (hope it's balanced):

2(NH2)2CO(s) + 7Ca(OCl)2(s) => 4NO2(g) + 2CO2(g) + 4H2O(g) + 7CaCl2(s) + heat

Bear in mind though that pool chlorine granules are only sixty something percent hypochlorite, 68% AFAIK?

Due to the potential violence of the reaction, I'd say passing the gasses directly through crushed ice is the best way to condense it as I think it would be evolved so rapidly that ordinary water-cooled condensers wouldn't cope with the required heat transfer rate. Also, be warned, this might explode (or do nothing or do nothing, then explode!).

Much of the carbon dioxide generated may well form calcium carbonate as commercial calcium hypochlorites probably contain significant amounts of lime AFAIK? Anyhow, goodness only knows what will really go on in this soup of a reaction (if anything).

[Edited on 20-6-2015 by deltaH]

aga - 19-6-2015 at 12:29

Nice one deltaH.

Great to see that i'ts not only me that posts when drunk ;)

deltaH - 19-6-2015 at 23:07

I'm hoping someone with urea can confirm if calcium hypochlorite is capable of oxidising it like I observed with melamine.

If that could work, it's game over for nitric acid production, well almost ;)

j_sum1 - 20-6-2015 at 00:36

deltaH, you always come up with good ideas. And it is good to see you back around here.
I have both urea and calcium hypochlorite in large quantities. As I recall, my Ca(OCl)2 is 100% which might be a bonus. I might just get some time to experiment in a week's time. I will grind up a 10g batch according to the stoichiometry you stated (after checking it first) and see what happens.

I rather suspect it will be a dud. But that is based on the notion that, if it was that easy, someone would have stumbled on this idea before. Still, that it what experimentation is for.

I will report back when done.

deltaH - 20-6-2015 at 01:21

Thanks j_sum1. Yeah, I'm back... now with 20% more danger and less referencing :D

Yes, it does seem too good to be true, but that said, melamine definitely reacts even when very diluted in water. It might just be that melamine is 'special' in this regard.

/speculation

I have a sneaky suspicion that the melamine-hypochlorite oxidation proceeds via formation of cationic amine oxide radical on the melamine's ring nitrogen position (stabilised by delocalisation), hence the extreme bright red colour and very high solubility of that intermediate (thinking of TEMPO here). i.e. something like this:

melamine oxide cationic radical.jpg - 20kB

I think when you add more oxygen to that molecule, it simply can't hold itself together anymore and starts to break up with the fragments oxidising fully.

/speculation off

So yeah, while urea's carbamate nitrogen groups would probably be far more inert, I think it's worth a shot since I did get the melamine-hypochlorite oxidation to work.


papaya - 20-6-2015 at 16:22

Where do you get melanine ?

battoussai114 - 20-6-2015 at 20:38

Quote: Originally posted by papaya  
Where do you get melanine ?

Darker skin and brown eyes mostly.

j_sum1 - 20-6-2015 at 20:59

Melamine chemical is different from melamine the popular name for the polymer although the two are related.
Melanin the pigmentation is completely unrelated. It is an etymological fluke that the two are only a couple of letters apart.

battoussai114 - 20-6-2015 at 21:08

Quote: Originally posted by j_sum1  
Melamine chemical is different from melamine the popular name for the polymer although the two are related.
Melanin the pigmentation is completely unrelated. It is an etymological fluke that the two are only a couple of letters apart.

Wait, in english the pigment is called melanin!? Dang it! Completely messed up :(

j_sum1 - 20-6-2015 at 21:10

Hey. I'm no expert. I just looked it up.
https://en.wikipedia.org/wiki/Melamine#Etymology

deltaH - 20-6-2015 at 22:24

Exactly so, melamine (the resin plastic that is used to make dishes) is made by reacting melamine powder (the chemical) with formaldehyde. This makes a very hard resin similar to Bakelite, except that instead of phenol units linked by methylene bridges, it's now melamine units linked by methylene bridges.

Melamine powder (the chemical) is not OTC as far as I know. I bought mine from a chemical supplier a very long time ago (more than ten years). I don't have any left, sadly :(

I did tons of experiments with my melamine powder and it's a pretty interesting chemical. For example, it forms nearly completely insoluble salts with virtually all mineral acids INCLUDING nitric acid! I don't know any other insoluble nitrate salt except for melamine nitrate, but I'm sure SM members will be able to name some ;)

If we get desperate (i.e. urea is not oxidizable by hypochlorite), then one possibility is to grind/file up a cheap melamine plate into a powder and try oxidizing that :o The extra methylene units will, however, consume much more hypochlorite per amount of NO2 you would make, which would be wasteful. I hope it won't come to such desperation!

j_sum, when you scope the initial test-tube amounts for the urea-hypochlorite reaction, if you get nothing initially, try gentle heating. I really don't want to have to file plates, besides, melamine resin might be too inert to react with hypochlorite.



deltaH - 26-6-2015 at 23:51

A cursory search of urea-hypochlorite is yielding surprising results. Many sites mention the formation of explosive nitrogen trichloride, which was unexpected as NCl3 requires acidic chlorinating conditions to form AFAIK.

It may be that the CO2 formed during reaction drops the pH by precipitating CaCO3.

While commercial calcium hypochlorite may contain lime as an impurity, very pure calcium hypochlorite might become more acidic as the reaction proceeds.

I strongly advise adding some lime to the material to ensure basic conditions and so minimise the risk of forming explosive NCl3 when working with very pure calcium hypochlorite.

It would also capture more of the CO2, thus decreasing evolved gas volumes.

The hypothetical equation would then be:

7Ca(ClO)2 + 2Ca(OH)2 + 2(NH2)2CO => 2CaCO3 + 7CaCl2 + 4NO2 + 6H2O

That said, having NO2 evolve under basic conditions is another suspect matter :mad: This reaction is proving more complex that I would have thought or liked :mad::mad:

If you calculate what the mass percentage of calcium hypochlorite is in the calcium hypochlorite + calcium hydroxide stoichiometry indicated above, it comes to 87%.

Thus pool 'HTH type' granular chlorine, for example, with specified 68% hypochlorite, may in fact already contain enough lime, assuming the residual mass is a mixture of calcium hydroxide and chloride.

This is what I was using with my melamine chemistry and may be why I didn't observe any formation of NCl3.

[Edited on 27-6-2015 by deltaH]

battoussai114 - 27-6-2015 at 19:23

I've been this close from trying to do thermal oxidation of urea due to the utter nothingness I found when looking for it online.... then it occurred that I don't have a fume hood and I'd probably be safer asking around here.

So, anyone have any idea on what I should expect to form from this sort of reaction? With the ease of getting refractory bricks it shouldn't be that hard to build a thermal oxidation reactor... heck, maybe some transition metal catalyst like those used in oxidation of other organic compounds could be used to lower the heat necessary.

deltaH - 27-6-2015 at 23:35

I don't see urea forming NOx by thermal pyrolysis because of the pathway that urea pyrolysis follows :(

There is a good thread on thermal decomposition of urea on here somewhere, use google's site search. There's also a nifty review paper on the topic on "cyanurates" that will give you ABSOLUTELY everything you want to know about it. If you don't find it, PM me.

Basically if you mix the appropriate amount of slaked lime and urea and heat, you will first produce copious amounts of ammonia :mad: and ultimately if you get to red heat, calcium cyanamide with the evolution of lots of CO2. Note: Without the lime, you will release toxic isocyanic acid in large amounts!

The calcium cyanamide can be converted to melamine in theory (by acidification with a calcium precipitating weak acid, eg. CO2 and boiling down the resulting cyanamide solution after filtering off the carbonate) though all in all, this is no trivial matter.

As I see it at the moment, there's two possibilities with urea:

(1) The hypothetical low-temperature oxidation with some super oxidant.
(2) Making ammonia from it and then catalytically oxidising the ammonia at high temperature with oxygen.

The second option is tough to execute, hence only option (1) seems plausible in practical terms, though whether the chemistry will play along is [highly] suspect at this stage.

[Edited on 28-6-2015 by deltaH]

battoussai114 - 29-6-2015 at 13:40

Quote: Originally posted by deltaH  
I don't see urea forming NOx by thermal pyrolysis because of the pathway that urea pyrolysis follows :(

There is a good thread on thermal decomposition of urea on here somewhere, use google's site search. There's also a nifty review paper on the topic on "cyanurates" that will give you ABSOLUTELY everything you want to know about it. If you don't find it, PM me.

Basically if you mix the appropriate amount of slaked lime and urea and heat, you will first produce copious amounts of ammonia :mad: and ultimately if you get to red heat, calcium cyanamide with the evolution of lots of CO2. Note: Without the lime, you will release toxic isocyanic acid in large amounts!

The calcium cyanamide can be converted to melamine in theory (by acidification with a calcium precipitating weak acid, eg. CO2 and boiling down the resulting cyanamide solution after filtering off the carbonate) though all in all, this is no trivial matter.

As I see it at the moment, there's two possibilities with urea:

(1) The hypothetical low-temperature oxidation with some super oxidant.
(2) Making ammonia from it and then catalytically oxidising the ammonia at high temperature with oxygen.

The second option is tough to execute, hence only option (1) seems plausible in practical terms, though whether the chemistry will play along is [highly] suspect at this stage.

[Edited on 28-6-2015 by deltaH]

Well, I'm really glad I decided to ask :)
Good readyng material from the urea pyrolysis topic, thanks for the heads up. Couldn't find the cyanurates paper but its okay since I don't have my proxy set up for getting access to RSC and others anyway.

I may start reading on catalytic oxidation of ammonia now that my dreams of oxidating urea have been crushed, but I really can't imagine a easily built reactor in this case since handling gases never worked well for me.... when I thought of the urea decomposition I had in mind a batch reactor that could be sealed with the solid compound inside and left to react without bothering with feed streams and stuff.... But oh well, its never that easy is it? :D

Lastly it probably should've occurred to me that pyrolysis and thermal oxidation were the obvious search keyword, not oxidation.

j_sum1 - 29-6-2015 at 19:20

Following deltaH's suggestion, I performed a couple of tests on urea and calcium hypochlorite. Here are the results.

Small scale: 0.1g urea, 0.13g Ca(OH)2 and 0.83g Ca(OCl)2 Drop of water to initiate.
A reasonable amount of gas evolution. Wishful thinking may have labelled it as brown.

Larger scale: 1.0g urea, 1.33g Ca(OH)2 and 8.33g Ca(OCl)2. Gas fed into H2O2. Drops of water via sep funnel to get the reaction going and speed it up if/when necessary.
Gas evolution began at low rate before water was added. At the first drop of water large volume of gas was produced. Lots of white fumes filled the flask and came through the apparatus. It was smokey in appearance and my guess is that it was particulate CaCO3. Reaction was very exothermic and flask became hot to touch.
System experienced suck-back which meant that I lost the H2O2. I added some more and continued to attempt to dissolve the gas. Again a wishful thought might claim a brownish tinge but it definitely was not pure NO2.
Solution at the end still contained a lot of peroxide (unsurprisingly). It was a very light brown colour. pH was around with indicator paper. No visible reaction between this and copper wire so I guess that HNO3 content is minimal.

Reaction with the Ca(OH)2: 0.2g urea, 1.7g Ca(OCl)2
Reaction began while powders were being stirred. Quite loud explosive pops which presumably came from NCl3. Rather an exciting experiment but ultimately a bit fruitless. I don't think I would want to upscale that at all. at 0.2g scale it was quite percussive on the glass beaker I was using.



So, interesting findings ,but probably no progress made towards the HNO3 goal.
Still, good to eliminate one possibility.

[Edit]
Adding photos
(1) Apparatus setup
(2) Gas evolution (after it had tamed a bit. I think the particulate suspension is CaCO3.
(3) Slight discoloration of the liquid but nothing remarkable
(4) No effect on copper wire
(5) But clearly acidic. Is this carboxylic acid?
Now the mission is to figure out what reaction has actually occurred.



[Edited on 30-6-2015 by j_sum1]

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deltaH - 30-6-2015 at 02:01

Thanks for experimenting j_sum1, I really appreciate the effort!

I'm glad my suggestion of lime was not a pointless precaution, in a moment of weakness, I thought I might have been silly suggesting it :D

You might be forming calcium nitrate in solution because of the necessary basic conditions.

Another idea is to swop the order of addition and try to target the formation of NO (nitric oxide) gas instead to avoid forming nitrates. This is then easily oxidised externally to NO2 gas.

e.g. Prepare a saturated urea solution and add the lime-hypochlorite incrementally/gradually --- weaker oxidising conditions targeting NO gas formation that should bubble out due to low solubility.

The reaction stoichiometry would change of course, but on the upside, you can economise on hypochlorite!

[Edited on 30-6-2015 by deltaH]

j_sum1 - 30-6-2015 at 03:36

You keep editing your post deltaH! :)

All good suggestions. I see merit in adding a solution to the solid. If nothing else it saves me the task of grinding both the urea prills and the hypochlorite. Before the next round however, I think I want to be a bit clearer on what the reaction products have been. There are obviously a few possibilities.

As for conserving Ca(OCl)2, it is the least of my worries. I have a couple of kilograms and more is available from the supermarket less than 300 metres from my house.

deltaH - 30-6-2015 at 04:30

Ha ha, yes I'm notorious for that, trying to decide what are good suggestions and then keep coming up with better trains of thought :P, so this is MkIV version lol

Very nice photos and work, thanks again!

j_sum1 - 30-6-2015 at 04:53

I like the one with the fuming test tube. I might have to do that one again just for the fun of it.

papaya - 30-6-2015 at 05:00

I'm sorry guys, but if I remember correctly in the old hydrazine thread this reaction (hypochlorite + urea in a big excess) was utilized to synthesize hydrazine hydrate, why do you thing you'll get other things out of it?

j_sum1 - 30-6-2015 at 05:03

Quote: Originally posted by papaya  
I'm sorry guys, but if I remember correctly in the old hydrazine thread this reaction (hypochlorite + urea in a big excess) was utilized to synthesize hydrazine hydrate, why do you thing you'll get other things out of it?

Because I don't know any better?
Seriously, playing with urea is new ground for me. I know next to nothing about it. I will have to look that one up.

deltaH - 30-6-2015 at 05:03

Because hydrazine is a reducing agent and there's excess very active oxidant? Not saying it's not forming, right now we haven't yet pinned down exactly what is forming in the oxidation :D

My big worry is that it's making N2 gas (and CO2). That would be rather useless aside for airbag generators, lol.

The oxidising conditions are very aggressive with calcium hypochlorite, but maybe hydrazine is forming in situ and then oxidising to N2?... which would be really bad :mad:

Way back, when mucking about with melamine, I found that sodium hypochlorite didn't oxidise melamine, but calcium hypochlorite did. So there's a difference there too, maybe.

Your first picture is very :cool: j_sum1

Maybe try making some coloured hydrazone's with your solutions to see if there's hydrazine in there? Do you have almond essence (not even sure if that makes a yellow hydrazone). In a test-tube, acidify with a few drops acid, add benzaldehyde solution and heat?

[Edited on 30-6-2015 by deltaH]

deltaH - 30-6-2015 at 11:59

Okay, so you probably want more proof that me 'just saying' that hydrazine shouldn't survive hypochlorite excesses. :P

There's a report from the US military at this link:

http://www.dtic.mil/dtic/tr/fulltext/u2/a213557.pdf

where they tested hypochlorites to decontaminate hydrazine spills. They found that the reaction of both liquid bleach and solid bleach (calcium hypochlorite?) on dilute hydrazine solutions was "rapid", highly exothermic and evolved almost entirely nitrogen gas for unsubstituted hydrazine.

So as I see it, if hydrazine was produced, then it would decompose very quickly to N2 in that reaction. If N2 is the main gas produced, then this is a possible reaction pathway: -NH2 => ClNH2 => NH2NH2 => N2

Else calcium nitrate may have formed if the gas evolved was not nitrogen (and carbon dioxide).

papaya - 30-6-2015 at 14:04

Yes, oxidation of hydrazine in most cases yields nitrogen gas, sometimes some amount of azide can be formed, which can also further be oxidized to nitrogen. IMHO what you should look for is a method (maybe some well known in organic synthesis) to oxidize amine (or amide) group to a nitro one, then maybe you have a chance to "extract" that in form of NO2 gas or nitrate.

j_sum1 - 1-7-2015 at 02:56

I fear that I have somewhat hijacked this thread and that my current pursuits are only tangentially related to the goal in mind (and are unlikely to get closer.)

With that disclaimer aside, I should report back on the calcium hypochlorite and urea adventures.
It seems that water is more than an inert reaction medium in this case. I added saturated Ca(OCl)2 solution to urea prills to observe what would happen. No Ca(OH)2 used this time.

I have to conclude that the presence of water changes the reaction that proceeds. I have really little clue as to the nature of the products.

The next thing to do is to attempt the reaction in dry conditions with the inclusion of Ca(OH)2. I might try some excess of the lime as a CO2 mop. I am going to predict minimal NO2 production if any. I will report back when done.

deltaH - 1-7-2015 at 03:35

Very interesting j_sum1.

Quote:
Immediate fizzing of the prills. Definite smell to the gas given off. Sort of chlorine-y but less sharp. Similar to the smell of most pool chemicals I guess, but maybe with a hint of ammonia to it. I am prepared to concede that there may be more than one gaseous species evolved.


Sounds like chloramine. Depending on the pH, you could be making chloramine, dichloramine and finally trichloramine the more acid the reaction gets.

It also sounds like the oxidation is too mild in solution to go beyond the 'standard' chlorinations of ammonia.

clearly_not_atara - 1-7-2015 at 08:46

In the presence of Br-, O3 takes NH3 >> NO3-:

http://www.sciencedirect.com/science/article/pii/00431354849...

Without Br-, O3 takes NH3 >> NO3- only at high pH. This reaction is mostly studied in wastewater treatment and has never been optimized to produce nitrate. However Br- is easy for us to come by and the mechanism sounds both plausible and likely to generate the right product (O3 reacts readily with NH2Br, it seems).

Ozonation is questionably OTC. Ozone is a toxic gas, after all. However its production is so well-studied there is even a tutorial reference for an undergraduate lab assignment, where sulfuric acid is electrolyzed:

http://pubs.acs.org/doi/abs/10.1021/ed082p1546
http://jes.ecsdl.org/content/129/3/506 (a comparison of electrolytic O3 generation methods)

And of course ozone generators are a well-known thing.

Quote:
I performed a couple of tests on urea and calcium hypochlorite.


I'm sorry, but I really don't think urea is going to work. An amide like urea is delocalized across the N-C-O bond system which forces the nitrogen into a planar configuration. Oxidations of amines usually proceed through tetrahedral intermediates, whereas amides undergo a reaction more like electrophilic aromatic substitution (amides react with TCCA or NBS but amines react with KMnO4 or K2CrO3) which usually produces a halo-amide. That rearranges to hydrazine, and once you're at hydrazine it's just too easy for an oxidation to produce N2 instead of the desired product. Best thing you can do with urea is probably to hydrolyse it.

[Edited on 1-7-2015 by clearly_not_atara]

deltaH - 1-7-2015 at 09:47

Interesting what you said about the configurations clearly_not_atara.

Ozone sounds like fun, but not the OTCishness we need.

But that got me thinking, if that works, maybe singlet oxygen might as well? That can be prepared from a hypochlorite and hydrogen peroxide according to Wikipedia, so it's OTC for some of us. If I understand you correctly, a bromide salt in there may act as a catalyst?

So I guess what I'm proposing is 'upgrading' the urea-hypochlorite reaction by adding hydrogen peroxide and a bromide salt, say KBr as catalyst? Is it really necessary to have bromine in the fray?

Hypothetical equation:

14H2O2(dil)+7Ca(ClO)2(s) + 4Ca(OH)2(s) + 4(NH2)2CO(aq) =>[Br-(cat)]=> 4CaCO3(s) + 7CaCl2(aq) + 8NO2(g) + 26H2O(l)

Exactly how one is supposed to excecute such an addition without blowing yourself up ... :o

Simultaneous dropwise addition of two solutions (H2O2 and separately, urea) to hypochlorite prills with some lime in a flask? :o:o The timing is going to be really hard to get right because of the very fast reaction of urea and hypochlorite alone. What a mess :mad:

Could be easier on a test tube scale with two pipettes and a simultaneous squart of a ml with a tiny amount of hypochlorite and lime in the bottom of the tube?

PS. Singlet oxygen is toxic AFAIK, be forewarned :o

[Edited on 1-7-2015 by deltaH]

battoussai114 - 5-7-2015 at 17:35

So I've been thinking of getting some copper vessels to test the catalytic activity of copper on ammonia oxidation (nothing more complex than space velocity and the effect of parameters such as temperature and residence time). But I'm not sure about how to test the conversion to NOx if ammonia is not completely removed. If all I got was N2 and NOx I could just run the gas through a wet scrubber containing hydrogen peroxide and then titrate to check how much acid was formed. But if there's also ammonia I'd get aqueous ammonia and then ammonia nitrate, and I'm not so sure how'd I check for it... sure I could just evaporate and weight what I got (assuming I got some deionized water and no contaminants in the H2O2), but that would be kinda problematic as some ammonia could pass through the scrubber without being dissolved, and I'm not too much into building a wet glass absorption column.
Also, any ideas on how to make a fluidized bed reactor without needing to pressurize the ammonia stream prior to running it through the reactor?

violet sin - 6-7-2015 at 01:37

"any ideas on how to make a fluidized bed reactor without needing to pressurize the ammonia stream prior to running it through the reactor?"

you need oxygen right? atmospheric O2 or canister O2 injection from the bottom, adjustable(needle valve) and constantly on during operation, with an interruptible and adjustable NH3 line that feeds into the O2 injection port. if NH3 is at minimal needed supply pressure, it shouldn't affect the bed too much, so it won't spray it out.

just a quick though

ave369 - 12-7-2015 at 02:41

My solution to the challenge, using what's available in my hometown.

Step 1: I go to the farmer market, buy a crocodile and a baby donkey.

Step 2: I kill the crocodile and the baby donkey.

Step 3: I stuff the crocodile with the baby donkey and put them in an old fuel drum.

Step 4: I start to heat the fuel drum with a campfire.

Step 5: I start to chant Hare Krishna and dance around the drum with a tambourine.

Step 6: I discard the tambourine, the drum, the crocodile and the donkey. I go to the spare car parts section of the town shopping arcade and buy battery acid. I concentrate the battery acid by boiling and evaporation of water, getting 96% H2SO4. I go to the gardening section of the town shopping arcade and buy any form of nitrate fertilizer. I put the acid and the fertilizer into a distillation apparatus and start to slowly heat the flask with mixture. I collect red fuming nitric acid. I dilute the red fuming nitric acid with water and start bubbling air through it until it becomes colorless.

Step 7: PROFIT

I am perfectly aware of the fact that this is boring, uninnovative and less fun than extracting nitric acid from crocodiles stuffed with donkeys. But it works.

[Edited on 12-7-2015 by ave369]

[Edited on 12-7-2015 by ave369]

j_sum1 - 12-7-2015 at 04:01

Interesting edit ave369. Your post is very different from when I first saw it.
Discussion early on in this competition clarified that use of nitrates from fertiliser was excluded as a method for this competition.

ave369 - 12-7-2015 at 05:29

As far as I know, the competition is no longer going on.

As for the edit... I've made it after I realized that using nitrate fertilizer is not kosher because it isn't interesting or innovative. So I made up a very interesting and innovative method involving crocodiles and donkeys. Shame it doesn't work, unlike the fertilizer method.

battoussai114 - 12-7-2015 at 07:55

Quote: Originally posted by ave369  
As far as I know, the competition is no longer going on.


As far as there is an "Open" beside the title of the discussion, its open :P

[Edited on 12-7-2015 by battoussai114]

aga - 12-7-2015 at 11:25

The Competition, and the Prize are still available.

Nobody has been able to reach the target.

The competition will Continue until someone wins the Prize.

I think i increased the prize from 250 euros to 350 euros at some point.

[Edited on 12-7-2015 by aga]

kecskesajt - 22-7-2015 at 03:41

Im thinking of the red hot copper wire method.Not wire,sponge(having acces to copper sponge :D) to minimalise the ammonia content in the end product.The ammonia generator would be NKP fertiliser(so not pure nitrate salt) and drain cleaner(unkown NaOH content, only says 40-60%).Heat the copper sponge in a copper pipe and run ammonia thrue it.
Can 35 % H2O2 used?
Water poorly absorbs nitrogen oxides.
I might post a pic about the apparatus.

Praxichys - 22-7-2015 at 08:08

Quote: Originally posted by kecskesajt  
Water poorly absorbs nitrogen oxides.


Try absorbing it in NaOH or Na2CO3. You will end up with sodium nitrate/nitrite in solution. Agitate it with air for a few days to convert the nitrite to nitrate.

j_sum1 - 22-7-2015 at 08:11

Yeah, but the goal is nitric acid and not merely nitrates. Otherwise you could simply take the fertiliser and -- you know, dissolve it in water.

SimpleChemist-238 - 22-7-2015 at 12:08

Easy, stump remover or cold pack, and PH down from publix. Distill and add water to get it to 30%.

Or

Calcium nitrate fertilizer and sulfuric acid. Then add water after filtration.

Texium - 22-7-2015 at 12:13

Quote: Originally posted by SimpleChemist-238  
Easy, stump remover or cold pack, and PH down from publix. Distill and add water to get it to 30%.

Or

Calcium nitrate fertilizer and sulfuric acid. Then add water after filtration.
Clearly you haven't read this thread. Take a few minutes and come back when you're done. Not so simple as that, SimpleChemist.

aga - 22-7-2015 at 12:16

sigh

Perhaps there should be a New DAC #5.5 Challenge :

50p for whoever can read the rules and totally ignore them the best, and who finds the FIRST and WELL KNOWN route to HNO3 the fastest.

Hey ! How about Distill a Nitrate salt with conc H2SO4.

Woohoo ! I won the DAC #5.5 !! Wheee.

Yawn.

aga - 22-7-2015 at 12:25

OK.

Many of you guys are much better chemists that i can ever hope to be.

Try to think more sideways rather than down the well-trodden routes that you have studied.

In other words, try to Use your knowledge imaginatively rather than spin around inside the safe, known areas.

kecskesajt is maybe onto something.

Perhaps the Copper need not be solid copper, but a complex with Silicon for
example.

I had the idea of using some nitrogen fixing plants in a hydroponic setup to get them to nitrate the water, but got lost when looking for what to do with it after that.

papaya - 23-7-2015 at 11:39

Find out this guy and send him money! (well, not nitric ACID, but nitrate)
www.youtube.com/watch?v=8wW5KR1pDxs

blogfast25 - 23-7-2015 at 12:12

Quote: Originally posted by papaya  
Find out this guy and send him money! (well, not nitric ACID, but nitrate)
www.youtube.com/watch?v=8wW5KR1pDxs


Hope you realise just how much organic material you need to process to get a modest yield of impure KNO3, you know! ;)

Keep your money and buy some stump remover.

[Edited on 23-7-2015 by blogfast25]

papaya - 23-7-2015 at 12:23

Yes, but you can start not from urine but urea, or maybe other organic nitrogen source say proteins. If one can isolate and grow a colony of that particular bacteria (hope it's only one) that's responsible for the process.. then nitrate becomes like brewing your moonshine!

aga - 23-7-2015 at 12:45

still in the 'normal' loop there, yet spinning off eccentrically.

Keep up the momentum and your mind may arrive in an area where you can find a solution.

Dangle89 - 1-8-2015 at 00:47

Holy flapping duckshit Batman! What a mission!

I started reading this whole thead a couple of nights ago and got sucked in by the drama /exitment of ideas floating around and experiments being conducted didnt end up getting to bed until about 3am :P Work sucked the next day :(

j_sum1 I admire your comitment to this project and I tip my hat to you sir :D Please keep going if you can as it is for a worthy cause :)

I wish I had some info / knowledge to assist the goal but I dont :-(

I would try myself but I get 1L of 70% HNO3 for about $10AU close to home and have never had to make it myself even from the simplest methods so "thinking outside the box", as Aga puts it, is beyond me here. Sorry guys :(

I wish you all the best of luck and look forward to reading more drama / excitment / breakthroughs and inevitably someone piping in with "H2SO4 + XNO3" :mad: shortly :D

Dangle

deltaH - 1-8-2015 at 01:25

Yes, it is certainly proving to be elusive. My money is still on the combustion of soybeans to form low concentrations of NOx in the flue gas.

papaya - 1-8-2015 at 05:56

deltaH, why to burn soybeans, because it contains some nitrogen ? Then why not to burn urotropin tablets - more nitrogen, cheap.. but it won't give you NO2, only N2 upon combustion. Why do you think burning proteins gives off NOx btw?

deltaH - 1-8-2015 at 06:07

I've already discussed this at length earlier in this thread. The production of NOx due to combustion has been heavily studied because of its contribution to air pollution. There are various mechanisms and sources for its formation, one of them, for example from diesel engine exhaust is because of the direct combination of N2 and O2 from air at high temperature, but there's a second known route, so-called fuel nitrogen, i.e. fuels that contain nitrogen in their structure. This is well known, see the discussion early in this thread. When you burn fuels that contain organic nitrogen, you pollute. I simply was going for the worst case scenario, i.e. use protein as fuel and the cheapest protein is soybeans.

The mechanism is complex, but it's critically important to have excess oxygen and as high temperatures as possible, else you will make a lot of N2. This is why this method would require a forced air type furnace to work effectively. I've even suggested a furnace with an internal propeller to smash the beans around to strip ash and maximise internal recycle for the excess air.

We've also discussed adding a catalyst, we settled on manganese dioxide (from soaking beans in KMnO4 solution, then drying) to help with the oxidation of the nitrogen an up selectivity for forming NOx versus N2.

[Edited on 1-8-2015 by deltaH]

papaya - 1-8-2015 at 06:25

Well, what I'm trying to say is that if you want to TEST that idea about high temp. combustion you could use urotropine, which is a well defined compound and not the horseshit source of nitrogen, at least in the beginning :)

deltaH - 1-8-2015 at 06:36

Hexamine fuel tablets was already suggested for a proof of concept test and I think it IS better suited IMHO.

Look, there's NO DOUBT that you can make NOx by burning nitrogen compounds, that IS the Haber process!!! The challenge is getting it to work precisely with horseshit/piss, cattle feed, etc. lol

[Edited on 1-8-2015 by deltaH]

Dangle89 - 2-8-2015 at 06:50

Hello all :)

I use to be a mechanic and it just occurred to me (cant remember if this idea what touched earlier - Sorry if it has as it is a very long thread!) catalytic converters in cars, trucks, whatever convert:

CO + HxCx (any form of left over hydrocarbon fuel)
= CO2 + H2O

And

NOx = N2 + O2

Would it work if you pulled the exhaust off BEFORE the catalytic converter and bubbled the exhaust fumes through water? Your water would probably form a lot more crap than HNO3! But I believe Aga said distillation is allowed :)


Just wanted to check if this idea had already been looked at and dismissed?

If not I will try find my old mechanic qualification books to see if it has the % of NOx from typical exhaust fumes before they hit the Cat converter to see if the amount of NOx is in a respectable enough quantity.

Also forgot to say it is illegal to start a car with out a cat converter here (Something stupid like a $40,000 fine) but I would be happy to try It in the name of science! :D

Dangle.

[Edited on 2-8-2015 by Dangle89]

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