| Pages:
1
2 |
Wolfram
Hazard to Others
Posts: 133
Registered: 13-10-2003
Member Is Offline
Mood: No Mood
|
|
NaNO2 failure
Airgun pellets were melted. (I assumed that they consisted of lead.)
The metal block formed was filed down with a file.
50g of the metal powder was slowly added in smal portions to melted 20g melted NaNO3. To my surprice the "lead powder" was actually floting
on the melted NaNO3. When the termprature raised even higher some brown salt looking like common rust began to form.
Could it be SnO? What did the pellets consist of, they were heavy and the metal was soft and melting at about 200-300.
(The reaction I was trying to perform was:
Pb + NaNO3 + heat -----> PbO + NaNO2)
|
|
|
chloric1
International Hazard
Posts: 1070
Registered: 8-10-2003
Location: GroupVII of the periodic table
Member Is Offline
Mood: Stoichiometrically Balanced
|
|
MAke sure you have good lead
Well that lead in those pellets might have been alloyed with antimony, tin, or ??? First I would try the lead sinker weights for fishing lines and if
you still can't get respectible yields go to McMaster Carr on the net and they have lead in any form you want.
Incidently, a few years back I feel in love with the potassium nitrate granulated sugar mix. I mixed 6parts KNO3 and 4 parts sugar and the smallest
glowing ember would send the comp instantly into a whitish purple flame and a hissing sound with grey smoke. LOVED IT! Anyways when it was all said
and done the salt had melted into a yellow mass with readily yielded Nitrogen Oxides with HCL and H2SO4. Problem is the purity may not be too great
but it would separate from the burnt carbon junk by dissolving in cold water.
Fellow molecular manipulator
|
|
|
Mephisto
Chemicus Diabolicus
Posts: 294
Registered: 24-8-2002
Location: Germany
Member Is Offline
Mood: swinging
|
|
Another way to nitrites / Just a thought
The reduction of an alkali nitrate with lead has the disadvantage, that the product is polluted with nitrate. And the separation of nitrate and
nitrite is (relatively) difficult. Although most chemists here haven’t a problem to buy pure nitrite, the unavailability of an OTC-method to
synthesise pure alkali nitrites is annoying.
My idea is to reduce a strong alkali nitrate solution with zinc (in presence of acetic acid). Than the zinc-ions can be removed with ammonia by
forming insoluble zinc hydroxide, which can be filtered of. The remaining acetate and ammonium-ions can be removed by heating (the alkali nitrite
shouldn’t be destroyed by this).
The product should be quite pure, but the zinc hydroxide precipitation might be difficult, because it is dependent to the pH of the solution (so too
much ammonia will dissolve the precipitation again).
|
|
|
Microtek
National Hazard
Posts: 827
Registered: 23-9-2002
Member Is Offline
Mood: No Mood
|
|
In my opinion, the most convenient method, at least in theory, is reduction with charcoal in the presence of Ca(OH)2. The only products are NaNO2,
NaNO3, CaCO3 and Ca(OH)2 so you can cycle your product through the process a few times and should then be left with only NaNO2.
I think mr anonymous posted about this method a while ago, and I have also tried it myself. I found that the molten nitrate/nitrite becomes so thick
when adding the Ca(OH)2/C mix that the melt puffs up and has a tendency to burn partially. The process can still be used, but I think it works better
if you don't add quite so much reducer in each cycle.
|
|
|
Wolfram
Hazard to Others
Posts: 133
Registered: 13-10-2003
Member Is Offline
Mood: No Mood
|
|
Could it possibly be..
Could it possibly be so that PbO2 or Pb3O4 formed? It was a reddish-brown unsoluble heavy junk.
|
|
|
unionised
International Hazard
Posts: 5102
Registered: 1-11-2003
Location: UK
Member Is Offline
Mood: No Mood
|
|
Microtek.
That's an expensive synthesis of washing soda. Sodium nitrite oxidises carbon (I just tried it).
|
|
|
Wolfram
Hazard to Others
Posts: 133
Registered: 13-10-2003
Member Is Offline
Mood: No Mood
|
|
Real powdered Pb still ..
Real powdered Pb still gives me redbrown-color salt instead of pale yellow wtf is this? 
|
|
|
Microtek
National Hazard
Posts: 827
Registered: 23-9-2002
Member Is Offline
Mood: No Mood
|
|
Unionized: I don't think so. The procedure comes from an old patent that is an improvement of an even older common method for producing nitrite.
The improvement consisted of using Ca(OH)2 instead of NaOH which means that the only solubles are the product and the unreacted reactants.
It is essential to use only enough carbon to reduce the nitrate to nitrite; that is one half mole of carbon per mole nitrate. The carbon is added
slowly, and diluted with Ca(OH)2 to the melted NaNO3.
|
|
|
unionised
International Hazard
Posts: 5102
Registered: 1-11-2003
Location: UK
Member Is Offline
Mood: No Mood
|
|
Cycling the product through the reduction stage repeatedly will reduce all of it.
With a mixture of nitrate and nitrite the nitrite will melt at a lower temp (OK I realise there will be a eutectic somewhere) and will diffuse through
the material and react with the charcoal before the nitrate.
I realise the method has been patented. The patent office don't check if a method works before they issue a patent.
The colour of lead's oxides will vary with the particle size and the crystal form as well as the oxidation state.
generally PbO exists as litharge (red), stable at room temp. On heating you can get massicot (yellow) and stable above 488C
Pb3O4 (red lead)
PbO2 (brown whenever I have seen it, but sometimes decribed as maroon) and also a high pressure form that is black.
There are also other oxides like Pb12O19, but that's just getting silly.
[Edited on 28-11-2003 by unionised]
|
|
|
Microtek
National Hazard
Posts: 827
Registered: 23-9-2002
Member Is Offline
Mood: No Mood
|
|
Why would the nitrite react before the nitrate ? As I see it, the method using carbon is no different from using lead or aluminum to reduce the
nitrate, except that much less carbon is needed.
Besides, as I said, the patent is just a variation on an established theme ( heating NaNO3 with alkali and carbon ).
In fact I think the Muspratt book which Polverone has on his website used this process for nitrite production.
|
|
|
Wolfram
Hazard to Others
Posts: 133
Registered: 13-10-2003
Member Is Offline
Mood: No Mood
|
|
Look at this:
Look at this:
http://bcis.pacificu.edu/~polverone/muspratt1/c-834.html
, it states that:
KNO3 + 2C ---> KNO2 + 2CO
So maybee one could use it for NaNO3 also..?
|
|
|
unionised
International Hazard
Posts: 5102
Registered: 1-11-2003
Location: UK
Member Is Offline
Mood: No Mood
|
|
Let me now when anyone gainsays the point that repeated reduction with an excess of carbon will reduce the nitrite too.
Generally, solids don't react well with one another because they can only mix at the points of contact. Liquids will react wth solids rather
better. First one to melt is the first one to react.
Sodium nitrite starts to decompose, even when it is on its own, when it is about 10 C above the mpt of NaNO3. (Merck)
That means you need good temperature controll. That's a bit hard to achieve given that the reactions are exothermic.
I'm not saying you can't get any NaNO2 this way. I'm saying that repeated reduction will ruin the yield unless you can separate out the
nitrite betwen runs.
|
|
|
Microtek
National Hazard
Posts: 827
Registered: 23-9-2002
Member Is Offline
Mood: No Mood
|
|
That is the point I've been making all along; you need a stochiometric amount of carbon, but since the addition of all the C+Ca(OH)2 in one batch
makes the melt unstirrable, I recommend adding perhaps one third of that ( in little portions ) to the melted NaNO3, then cooling, dissolving in
water, filtering off the CaCO3 and any Ca(OH)2 and carbon, evaporating the water and then do the same with the rest of the C+Ca(OH)2.
|
|
|
unionised
International Hazard
Posts: 5102
Registered: 1-11-2003
Location: UK
Member Is Offline
Mood: No Mood
|
|
And the thermolysis of the nitrite?
What guarantee is there that, having been freshly generated in the presence of the hot charcoal, the nitrite will (because it has read a patent) defer
to the (possibly still solid) nitrate, get out of the way, and let it take its turn?
(Merck gives the decomp temps for the K salts and the nitrite is easier to decompose this tends to indicate that it will react quicker too)
|
|
|
Microtek
National Hazard
Posts: 827
Registered: 23-9-2002
Member Is Offline
Mood: No Mood
|
|
If that was the case, the nitrite would be decomposed by all of the reactions which involve the melting of a nitrate. I think the rate of thermal
decomposition is actually quite low for both KNO3 and KNO2. This is based on my attempts at producing nitrite by the supposedly common laboratory
practice of heating a nitrate. I heated it strongly; way above the melting point but gas evolution was very slow.
|
|
|
Wolfram
Hazard to Others
Posts: 133
Registered: 13-10-2003
Member Is Offline
Mood: No Mood
|
|
So go and try.
So please someone go and try:
NaNO3 + 2C ----> NaNO2 + 2 CO
use active carbon not barbecue cole. 
..come back and report the results.
What is the meaning with involving
Ca(OH)2?
If anyone doesn´t try it I will, but it will take some time becouse I have much to do and have no activated carbon.
[Edited on 1-12-2003 by Wolfram]
|
|
|
unionised
International Hazard
Posts: 5102
Registered: 1-11-2003
Location: UK
Member Is Offline
Mood: No Mood
|
|
I know it' not the same thing but I did the experiment with charcoal (activated as it happens) and NaNO2.
It flashed in much the same way as C and NaNO3.
That's why one of my earlier posts says "I just tried it"
The question people now need to address is why would this reaction not go in the presence of hot NaNO3.
Please let me know.
|
|
|
Microtek
National Hazard
Posts: 827
Registered: 23-9-2002
Member Is Offline
Mood: No Mood
|
|
I have tried it. With NaNO3, ordinary charcoal from wood, and Ca(OH)2. I think the reason that alkali is inkluded is to moderate the reaction so it
doesn't flash, but decomposes in a controlled manner. The product evolves nitrous fumes when reacted with acid, but I haven't done
any quantitative experiments to determine the amount of nitrite in it.
|
|
|
Wolfram
Hazard to Others
Posts: 133
Registered: 13-10-2003
Member Is Offline
Mood: No Mood
|
|
Wtf
Wtf in chemistry you become surpriced almost every day. I have thought that PbO is yellow. Now I read that there are two forms of PbO; yellow AND
red!
This explains why the Pb powder got red-brown in my experiment.
[Edited on 2-12-2003 by Wolfram]
|
|
|
JDP
Harmless
Posts: 43
Registered: 31-10-2003
Location: CA
Member Is Offline
Mood: |:booM
|
|
I believe the red vertion is Lead Tetroxide(Pb3O4). I've done the reaction of NaNO3 + Pb + heat be for and I got a fine tan to white pewdered
PbO.
~JDP
|
|
|
unionised
International Hazard
Posts: 5102
Registered: 1-11-2003
Location: UK
Member Is Offline
Mood: No Mood
|
|
How does the NO3- or NO2- know that the Ca(OH)2 is present? (it's barely soluble in the melt.)
If it is un-aware of this it will react as I observed.
Would anyone who wouldst reply to this please address this and my other questions rather than saying "its in a patent" or "it makes a
difference", without explaining why.
|
|
|
Microtek
National Hazard
Posts: 827
Registered: 23-9-2002
Member Is Offline
Mood: No Mood
|
|
One of the essential principles in science, is that no amount of theorizing can undo the findings of an experiment.
So, if an experiment has been conducted ( and I have conducted one ) that shows that it works, then that isn't changed by whether you or I can
understand how. Besides you haven't offered any solid evidence that the nitrite "should" be reduced before the nitrate either.
Saying that nitrite has a lower decomposition temperature than nitrate is not sufficient to conclude that it must be reduced faster. As I have said
before, if nitrite decomposed faster than nitrate, it wouldn't be possible to produce it by heating NaNO3.
As for the mechanism of moderation by Ca(OH)2, I can offer one which is nothing more than a plausible guess:
I you have a melt of NaNO3 and you pour powdered charcoal into it, the C will be oxidized by the nitrate just as in blackpowder. And just as in
blackpowder, the heat of reaction will accelerate the oxidation beyond control.
Now, if you add something like Ca(OH)2, the released energy will be absorbed by the inert ( in this context ) matter and so reduce the reaction rate
to a controllable level.
Whether this is the actual mechanism or not ( and I doubt that it is ) it demonstrates that additives doesn't have to be in solution to affect
the reaction.
|
|
|
TrollchEmist
Harmless
Posts: 5
Registered: 4-12-2003
Member Is Offline
Mood: No Mood
|
|
Did this reaction:
NaNO3 + C ----> NaNO2 + CO
Powdered NaNO3 and stoichiometric amount of active charcoal dust was mixed and then melted. Mixture was black-grey coloured. When melted some bubbles
evolved and then one spark(then I moved bit further and watched the fun part.) Huge blaze evolved from mixture. After reaction mixture colour was
changed to pale yellow-white. There was only minimal amount of charcoal left.
Seems a good produce for NaNO2
-TrollchEmist-
|
|
|
Wolfram
Hazard to Others
Posts: 133
Registered: 13-10-2003
Member Is Offline
Mood: No Mood
|
|
Congratulations!!!
Congratulations!!! 
Did you really set it on fire? How long did the fire last?
If you would have a good scale you could messure if the corect mass has become gas...
But you can instead buy a rat and check the effect. The LD dose for a human (75Kg) is about 3g. So for a 500 g rat it should be 3/150=0.02 g.
..no im joking but if would be nice to know if its the nitrite you have.
[Edited on 4-12-2003 by Wolfram] 
[Edited on 4-12-2003 by Wolfram]
|
|
|
TrollchEmist
Harmless
Posts: 5
Registered: 4-12-2003
Member Is Offline
Mood: No Mood
|
|
Yes it blazed really furiously about 10s or something(it was beautiful flame(height about 30cm and bright).
here the calculations:
58g NaNO3 and about 8g C
that would make a 0,68mole reaction
then (after filtering) getting NaNO2 powder(very white with pale yellow) about 31g. that's little bit less than expected but some of the blazing
mixture propably splashed away( there was amounts of that around heating place.
Then tested the pH of my 'product' in water and that was definitely base(pH paper)
The NaNO2 solution in water is weak base
Then I did little sulfuric acid test with my 'product' and NaNO3 side by side. (dropped little both of them into sulfuric acid) The
'product' evolved some gases(nitrogen oxides) while NaNO3 was still. After couple minutes 'product' was dissolved and NaNO3 was
still.
Broked my thermometer so i haven't done melting point test.
'Product' is definitely NaNO2
This is unbelievable easy synthesis of NaNO2
[Edited on 4-12-2003 by TrollchEmist]
|
|
|
| Pages:
1
2 |
![[*]](images/xpblue/default_icon.gif){kind=icon}
