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mayko
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[*] posted on 30-8-2013 at 17:46
A Thread of Negative Results


In order to prevent publication bias, I suggest a thread in which we post things that we tried that didn't work.

____________________________________


I tried to make nitrotobacco, on the theory that the cellulose in the plant matter could be nitrated similarly to cotton. (I am living in a smoking household and have a bit of frustration...) In retrospect, it was unlikely to work; cotton is nearly pure cellulose whereas dry tobacco can contain as much as 8% nicotine and ~20% minerals, according to Wagner's Chemical Technology.

Anyway, I used a recipe I'd used before, 2 parts sulfuric acid to 1 part nitric. I added a few grams of grape-flavored cigarillo tobacco to the mixture.

Long story short, a pretty violent and smelly nitration took place, and the shredded tobacco was reduced to a sludge.

Incidentally, Wagner's notes that potassium nitrate is found in tobacco, but doesn't influence combustion. “The combustibility of tobacco is not proportional to the quantity of nitric acid (nitrates) it contains”, citing two counterexamples.


[Edited on 31-8-2013 by mayko]




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elementcollector1
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[*] posted on 30-8-2013 at 17:54


I'd fill up this thread pretty quickly.
My most recent failure was attempting to isolate some CeCl3 from its hydroxide. The hydroxide was tan, the HCl solution clear. It turned yellow, indicative of some other anion or impurity (as CeCl3 crystals are perfectly clear, I surmise a solution is as well). I'll still try to crystallize it, but I don't have high hopes.




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[*] posted on 31-8-2013 at 06:07


just had one today, i have yet to test its burn to see if its all fucked up..

so.. nitration of cellulose
approx 10 hours, no problem
decent washing
then to boiling it in NaHCO3
solution turns darker
and darker..
i decide to try and change NaHCO3 for dragging out even more impurities
i realise the NC is all....
porous..
only small amounts i have of it to test
seems low nitrated and leaves behind residue, but how did i manage to remove NO3 from NC?? this is an industrially used method!!




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plante1999
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[*] posted on 31-8-2013 at 06:44


In industry, nitrocellulose is boiled in extremely dilute sodium carbonate solution, a better way for home chemist would be to boil in 2-3% urea sol..



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[*] posted on 31-8-2013 at 20:42


Tried making sodium by electrolysis of molten sodium hydroxide, it failed because I used a metal container, and the electricity just went through the container instead of the liquid >:-|

Also for some reason, the sodium hydroxide just disappeared after I heated it, maybe it boiled away?
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elementcollector1
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[*] posted on 31-8-2013 at 21:30


Quote: Originally posted by Cou  
Tried making sodium by electrolysis of molten sodium hydroxide, it failed because I used a metal container, and the electricity just went through the container instead of the liquid >:-|

Also for some reason, the sodium hydroxide just disappeared after I heated it, maybe it boiled away?


If you mean it lost volume, molten solids tend to do that - the only exceptions I know of being water and antimony.

Additionally, a metal container shouldn't affect things much - think of it as simply a larger electrode. Besides, metal can't contain electricity - it would simply spread it to something else, in this case back to the sodium ions. Was the container itself affected at the end of the run?




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[*] posted on 31-8-2013 at 21:34


It was stainless steel and I think the sodium hydroxide might have reacted with it. But what's really weird is, when I heated the sodium hydroxide, it started bubbling and making fumes, and it definitely wasn't hot enough to boil. So it could be that it was hydrated. BUT, if I keep heating it, the molten stuff just DISAPPEARS INTO THIN AIR, and all that's left is nothing inside the container. If I stop heating it, it leaves behind water and nothing else. Something really weird is happening, maybe I should post a video.

I'm going to try it again once I get a porcelain or clay crucible instead of metal, and use a wider volume so the sodium doesn't travel to the other electrode and explode. Is a 5 volt 1 amp power supply good or is that too much?
[Edited on 1-9-2013 by Cou]

[Edited on 1-9-2013 by Cou]

[Edited on 1-9-2013 by Cou]
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elementcollector1
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[*] posted on 31-8-2013 at 21:40


Quote: Originally posted by Cou  
Perhaps the sodium hydroxide is dissolved in the water that's left?


That's the guess I would make, if anything. Stainless steel should hold up fine to molten NaOH - unless you saw a trace of yellow or green color in your melt. Then there's a problem.
What's your heat source? NaOH does indeed boil at 1388 C, so unless you're using a blacksmith's forge or some such, that's not possible. Rather, I would think that the water in the sodium hydroxide (because it's really hard to make it anhydrous) is dissolving the NaOH as it heats (more heat, more solubility) and subsequently boiling off. Doesn't explain where the NaOH itself went, though... Sure you don't have a leak or unseen hole somewhere?




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[*] posted on 31-8-2013 at 21:42


I don't know what happened but I'm trying again once I get money for an order from elementalscientific.net (i hate being under 16, can't get a job, not enough money for my hobbies; now I see why most people that do this are adults)

[Edited on 1-9-2013 by Cou]
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[*] posted on 31-8-2013 at 23:29


Ha, were both in the same situation, saving up money for an Elemental order. Harder to save when youre under 16.



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[*] posted on 1-9-2013 at 05:01


well plante, i did use a very strong solution, it was alot of NaHCO3 i filled in
i cant really get urea (i know..)
i havent seen it anywhere in any garden shops, ever. and thats a long time..
perhaps it was so hot i managed to decompose the NC in situ, hossam have confirmed that the solution should darken, but im still confused over how my resulting NC turned to be..
actually doing some tests before NaHCO3 it did leave abit of residue, but i still think that it leaves alot more residue after the NaHCO3.. seemingly it was poorly nitrated so i might have another go on this, perhaps it became that consistency because of how poorly nitrated it was




~25 drops = 1mL @dH2O viscocity - STP
Truth is ever growing - but without context theres barely any such.

https://en.wikipedia.org/wiki/Solubility_table
http://www.trimen.pl/witek/calculators/stezenia.html
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[*] posted on 1-9-2013 at 06:04


Quote: Originally posted by elementcollector1  
I'd fill up this thread pretty quickly.
My most recent failure was attempting to isolate some CeCl3 from its hydroxide. The hydroxide was tan, the HCl solution clear. It turned yellow, indicative of some other anion or impurity (as CeCl3 crystals are perfectly clear, I surmise a solution is as well). I'll still try to crystallize it, but I don't have high hopes.


As so often, Fe3+ is a likely culprit. You need Fe free HCl and that's something you might not be able to get from Ye Olde Hardware Store (although I got some from there that was pretty colourless).

To separate the Ce and the Fe, try this. Acidify the solution and saturate the solution of CeCl3 with K2SO4 (use Wiki solubility table data for quantities to be used) by stirring in this salt directly, then simmer for 1/2 hour (or so). The Ce will precipitate as a white double salt: Ce2(SO4)3.K2SO4.3H2O (I'm actually not sure it's a trihydrate or a dehydrate but it matters not one iota here). The iron remains in solution, provided pH is lower than 4.

Allow to cool and chill to further reduce solubility of the Ce double sulphate, then filter. Rinse filter cake with cold, acidified saturated K2SO4. The filter cake should now be entirely free from iron.

Convert the double salt back to Ce(OH)3 with strong ammonia and filter and wash to get rid of K2SO4.

[Edited on 1-9-2013 by blogfast25]




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elementcollector1
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[*] posted on 1-9-2013 at 08:43


Quote: Originally posted by blogfast25  


As so often, Fe3+ is a likely culprit. You need Fe free HCl and that's something you might not be able to get from Ye Olde Hardware Store (although I got some from there that was pretty colourless).

To separate the Ce and the Fe, try this. Acidify the solution and saturate the solution of CeCl3 with K2SO4 (use Wiki solubility table data for quantities to be used) by stirring in this salt directly, then simmer for 1/2 hour (or so). The Ce will precipitate as a white double salt: Ce2(SO4)3.K2SO4.3H2O (I'm actually not sure it's a trihydrate or a dehydrate but it matters not one iota here). The iron remains in solution, provided pH is lower than 4.

Allow to cool and chill to further reduce solubility of the Ce double sulphate, then filter. Rinse filter cake with cold, acidified saturated K2SO4. The filter cake should now be entirely free from iron.

Convert the double salt back to Ce(OH)3 with strong ammonia and filter and wash to get rid of K2SO4.



This is odd, because the HCl was transparent and colorless. Nonetheless, I''ll give it a go when I get back, assuming the CeCl3 crystals (they should be evaporating out) contain some iron in them.

Does the double salt convert by itself with strong ammonia (Ce2(SO4)3.K2SO4.3H2O + NH4OH -> (NH4)2SO4 + Ce(OH)3 + K2SO[sub4

Or is heating required?

[Edited on 1-9-2013 by elementcollector1]




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[*] posted on 1-9-2013 at 09:36


EC1:

No, no heating required, just a bit of time (a few minutes).




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[*] posted on 1-9-2013 at 10:34


This thread seems like it's becoming a confusing mess. Although I like the idea of reporting negative results, I'm not sure that the execution is feasible—at least not in this manner.



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mayko
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[*] posted on 1-9-2013 at 11:16


The topic I had in mind was closer to "Things which appear to be infeasible", rather than "Things I am having difficulty with", which is what the general forum is for. I'm glad it generated interest and discussion though!



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[*] posted on 1-9-2013 at 12:40


It's certainly by its very nature a topic for hopelessly... going off topic!



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[*] posted on 2-9-2013 at 00:36


There is one good thing about this topic and that is that it invites people to post things which they otherwise most likely would not post. Posting failed attempts may be as useful as posting success stories. Others may learn from it and you yourself may learn from it as well, simply because you think over the whole process while writing down the results.



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[*] posted on 2-10-2013 at 03:52


I put eletrolised CuSO4 but nothing happened. what the hell?
i also added cotton in a nitrating bath nitric/sulfuric acid. the cotton dissolved, the solution turned green, then an orange percipitate formed.




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[*] posted on 2-10-2013 at 04:58


What exactly did you do with the CuSO4? Be more specific and describe your experiment in much more detail if you want a useful response.

What most likely happened in the cotton-experiment is that the cotton is broken down completely. The green color most likely is due to formation of NO2 and NO, which together form a blue compound N2O3, which in combination with red/brown NO2 gives a green color.
The orange precipitate may be polymeric/condensed and somewhat carbonized crap.

Try the experiment with clean white wadding. Add this to a 1 to 1 (by volume) mix of conc. H2SO4 and 65% HNO3.




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[*] posted on 2-10-2013 at 06:23


Quote:
I put eletrolised CuSO4 but nothing happened. what the hell?

On the assumption that you tried plating out copper to leave dilute H2SO4, your anode may have been unsuitable, or your PSU wasn't up to the job . . .
Your other problem would be that you used synthetic cotton!
Pharmacies for pure cotton . . . ?




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[*] posted on 2-10-2013 at 09:23


I used a carbon electrode and 12v battery like in NurdRage's video.
http://www.youtube.com/watch?v=5dUSF9Gl0xE
i also tried to make an ammonia solution, but the tubing was insucure and it came of mid reaction and sprayed ammonium gas in my face. (that was horrid).




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[*] posted on 21-10-2013 at 16:57


Quote: Originally posted by elementcollector1  

If you mean it lost volume, molten solids tend to do that - the only exceptions I know of being water and antimony.


I have quite a bit of gallium and I have noticed this as well




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[*] posted on 25-11-2013 at 17:55


Before I figured out what I was doing, I tried making black powder (Job related--aren't I lucky?). I thought that if charcoal was good,
then the highest purity 99.9% graphite flake would be even better. So donning face shields, leather aprons, and heavy gloves,
we tried igniting a bit of it (unconfined) with the burning end of a long stick. It was then that I learned firsthand just how hard it is to
ignite graphite, even with an oxidizer mixed with it. The paper just burned out from under it and left a little pile of my
"black powder" sitting there on the ground.
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[*] posted on 25-11-2013 at 19:18


Trying to make zinc sulfide by reaction of sulfur with zinc powder. I used the stoiciometric amounts of each compound, when I mixed them up in a mortar and tried to ignite it with an ethanol torch the reaction just didn't wanted to start. The sulfur and the zinc were lab grade, also I provided a good mix of both compounds so I think that it could have been the fault of the ethanol torch. Was the fault of the ethanol torch even considering the instability of that mix?
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