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Σldritch
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Quote: Originally posted by Melgar  |
Blue Fe(II)SO4 mixed with yellow sulfur would give a green color. Eventually, the iron II sulfate would oxidize to iron III, which is more of a
yellow/orange/brown color depending on concentration. |
I dubt that is what caused the color because it was very dark green to the point that after i washed out the beaker and neutralized some sodium
bisulfate in it the color returned albeit not nearly as dark and it seemed to move down to the bottom of the solution.
Anyway the aqeous reduction with liver of sulfur was a failure as expected, worth a try atleast.
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theAngryLittleBunny
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I tried reducing KNO3 with tin today, and I think I should tell you what happend and why it isn't a good idea.
So I melted 62g of KNO3 in a metal can on a gas burner and put a 35g piece of tin into it. It obviously quickly melted and on stirring, within a few
minutes a lot of gray SnO appeared and it seemed to be going well, the mixture became quite thick from the SnO. But after a substantial amount of SnO
built up, the KNO3/KNO2 started reacting with it in a thermite like way, a ton of SnO2 smoke was released from the steel can, and as soon as I saw
that the can started to melt, I knocked it from the gas burner and it cuntinued reacting violently for a like 30 seconds and melted completely through
the container.
So just don't use tin for this, it's not a good idea.
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Σldritch
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Very OTC Sodium NItrite
I found a really nice nitrite synth on atomistry:
The Sodium nitrite, NaNO2, is obtained by reducing sodium nitrate with metals such as lead or iron, with sulphur or carbon, or with material
containing these substances. In Dittrich's process the nitrate is heated with slaked lime and sawdust, the yield being almost quantitative, where as
the action of coal and charcoal is too energetic:
Im very suprised this has not been posted before. I have read about people here attempt something similar with charcoal but getting bad results. It
seems any reducing agent with a base will reduce nitrate to nitrite.
I did not have any sawdust nor calcium hydroxide on hand so i tried it with flour and sodium hydroxide drain cleaner:
12 NaNO3 + 12 NaOH + C6H10O5 = 12 NaNO2 + 6 Na2CO3 + 11 H2O
I heated 115g of Sodium Nitrate with the stochiometric ammounts of the other reagents, mixed sloppily together in a steel can with plenty of head
room, on a burner. A lid was loose placed over it. As it heated up it started smoking and foaming. The reaction seemed to stop after the smoking did
and i was left with a light yellow mixture at the bottom of the can.
Afterthe cake had cooled it was crushed up and dissolved in boiland water and cooled in a fridge. The mixture froze to a slush. It then filtered it on
a vacuum pump until mostr of the slush had melted. (The slush was probely a mixture of sodium carbonate and some sodium nitrite hydrate). The mixture
was then boiled down to the theoretical volume of a satured solution containing the theoretical amount of sodium nitrite formed in the reaction and
chilled and filtered again.
Then i thought i would neutralize the residual sodium hydroxide so i added sodium bicarbonate repeated the filtering again. I really should have
thought about that earlier but oh well.
I evaporated the final solution on a steam bath and tested with hydrochloric acid. It seems i got a relativly high yield even with extra steps though
it would probely be pointless to weigh.
This seems WAY better than using lead or sulfur or some other obscure ways to reduce nitrite. And with calcium hydroxide you could make it even
simpler. Nitrite can be made dirt cheap with this method im sure.
I really recommend this, im going to make some isopropyl nitrite now aga 
[Edited on 29-9-2017 by Σldritch]
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UC235
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http://www.sciencemadness.org/talk/viewthread.php?tid=52
Separation of unreacted nitrate is very difficult, and some brown gas on acid addition is hardly a good way to quantify such a mixture. Converting the
crude mix to a water-immiscible nitrite ester is probably among the best approaches.
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clearly_not_atara
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This is a pretty cool idea. Not sure why lime would be simpler... the insolubility of resulting CaCO3 might make it difficult to extract the product.
Solubility at 0 C:
KNO2: 280 g / 100 mL
KNO3: 13 g / 100 mL
This is large enough that a saturated solution prepared from the potassium salts at 0 C should have a sufficiently high proportion of nitrite for
practical purposes.
If nitrite free of nitrate is desired, such as for the preparation of N2O3, this reaction can be used:
NiCl2 (aq) + 6KNO2 (aq) >> K4Ni(NO2)6*H2O (s) + 2KCl (aq)
"Potassium nitrite (80 g in 25 ml water) was addedwith brisk stirring to NiCl2*6H20 (20 g in 20 ml water). The crystalline precipitate was
filtered, washed with cold methanol, and dried in the air (yield 88%)"
http://www.publish.csiro.au/CH/CH9731663
Potassium hexanitronickelate monohydrate precipitates as orange-brown crystals which dry to a violet solid when heated under vacuum. However, I am not
sure if this solid can be used to generate N2O3 by rxn with acids (although I suspect the answer is "yes" so long as the acid is strong enough, eg
H2SO4).
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Σldritch
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Misremembered the solubility of sodium nitrate... Anyway if the reaction was not complete i would not expect a white cake after the reaction but a
grey one. Of course the nitrite might react faster than the nitrate with the carbon so it does not guarantee purity but nitrate as the stronger
oxidizer should react first.
Ill try making a alkyl nitrite soon but i have guests now so it will have to wait.
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brubei
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Sodium Nitrate is commonly sold for ceramic making.
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Sulaiman
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OTC sodium nitrite is legal and easy in UK http://www.ebay.co.uk/itm/1kg-Sodium-nitrite-high-quality-/3...
do I get a prize ? 
CAUTION : Hobby Chemist, not Professional or even Amateur
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XeonTheMGPony
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Only if he'll ship to Canada!
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Pulverulescent
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| Quote: | I found a really nice nitrite synth on atomistry:
The Sodium nitrite, NaNO2, is obtained by reducing sodium nitrate with metals such as lead or iron, with sulphur or carbon, or with material
containing these substances |
Talking about ebay, you could exploit your easy prep. there ─ and, er, cash-in? 
"I know not with what weapons World War III will be fought, but World War IV will be fought with sticks and stones"
A Einstein
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symboom
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NiCl2 (aq) + 6KNO2 (aq) >> K4Ni(NO2)6*H2O (s) + 2KCl (aq)
Potassium hexanitronickelate
Interesting I wonder if using sodium nitrate with kcl would work
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clearly_not_atara
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Just use the potassium
salts, it's much easier.
In particular, K2HPO4 is much more soluble (140% w/w) than Na2HPO4 (7% w/w), so a saturated solution of the former should precipitate the latter.
KHCO3 has this property to a lesser extent.
Salt metathesis is highly underrated, I see.
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Σldritch
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I have a hard time getting potassium salts. The best i can get is 40%KCl 50%NaCl 10%MgSO4 mineral salt. I have had no sucess separating the KCl. I
have a little bit of KNO3 left bought from a now closed down store though. Id rather use a small excess of flour.
Im pretty confident that there is not much nitrate in it because if there was the nitrite would have to react faster than the nitrite in the molten
mixture. If that was the case almost no nitrite would be formed at all. It makes sense too since nitrate is a stronger oxidizer in these conditions.
(Ex. permanganate and ferrate)
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clearly_not_atara
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They don't sell cream of tartar where you live? Burning this salt gives potash.
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The Volatile Chemist
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cream of tartar is expensive OTC and burning it would reduce it to even less mass. Not an economical method.
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clearly_not_atara
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By this method you would be paying about $35/kg for potash with Amazon prices. It's not cheap but considering the costs associated with performing any
kind of amateur chemistry it doesn't sound expensive unless you're making silly amounts of nitrites (and for what?). OTOH K2SO4 is available on Amazon
for less than $10/kg but maybe he can't buy that.
I'm a bit surprised you couldn't just buy some kind of potassium fertilizer, given it's one of the big three plant minerals. Or you could always do
this:
NaNO3 (aq) + KCl (aq) [0 C or lower] >> KNO3 (s) + NaCl (aq)
OP apparently has access to KCl salt replacement so this should work, although it requires redoing the whole process.
[Edited on 1-10-2017 by clearly_not_atara]
[Edited on 1-10-2017 by clearly_not_atara]
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j_sum1
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I'm with Sulaiman on this one. In my world, sodium nitrite is OTC. And easier to obtain than nitrates.
http://www.ebay.com.au/itm/100g-bag-of-Sodium-nitrite-100-Fo...
http://www.melbournefooddepot.com/buy/sodium-nitrite-powder-...
[edit] typo
[Edited on 2-10-2017 by j_sum1]
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Σldritch
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What is OTC varies a lot
I would call something OTC when i can buy without giving my credit card imformation. That includes most stores where i live.
I can buy sodium nitrates in about half of the grocery stores here yet potassium salts are hard to obtain. I think this is intentional as terrorism
preventation, i doubt it is very effective though. I have not found any stores that sell sodium nitrite.
Anyway, i dont have a lot of calcium hydroxide which you would need do do this with potassium because of the similar solubility of potassium nitrite
and potassium carbonate. If you do not use calcium hydroxide it is more of a tradeoff between carbonate and nitrate impurities.
Mineral salt would probably introduce more impurities than it would help reduce.
Maybe i can titrate it? I tried permangante and ammonium chloride, neither seemed to work very well.
Also i enjoy the challenge of obtaining chemicals i can not buy.
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clearly_not_atara
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You wouldn't find potassium fertilizer in a grocery store. (Neither would I) You have to go somewhere that sells fertilizer. It'd be at a hardware
store for me -- the same place that sells tools and paint (Home Depot). They have giant bags marked "Potassium Sulfate" with a K2O percentage marked.
You would also see it at a garden supply.
It's really strange to ban potassium because it makes a pretty crappy bomb. It doesn't react and it's heavy which reduces the temperature of mixtures
containing it. And it's extremely useful as fertilizer. It's one of the most legitimate chemicals I can think of -- everyone from subsistence farmers
to yuppies washing their faces with Dr. Bronner's uses potassium.
But it tastes bitter and it's not a common ingredient in food, so you won't usually see it at the store.
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Broken Gears
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Sodium Nitrite should be easy to find OTC, as it's used in preserving meat.
Any Home-produktion, hunters/butchers shop or DIY beef jerky shop should have it OTC.
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karlos³
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Isn´t the stuff used for preservation of meat, "curing salt", just like 1% NaNO2 at most, the remainder being NaCl, NaNO3 for the major part?
At least in europe the curing salt contains per law less than one percent, usually.
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Melgar
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Potassium is just less common than sodium in general, and the salts are correspondingly less common. Interestingly though, plants contain almost no
sodium at all, which is why we prefer the taste of sodium salt on our food, and why deer are attracted to salt licks, etc. Humans and other animals
NEED sodium in our diets. Plants were able to evolve to not need sodium at all, and instead are able to extract potassium from feldspar (the most
common mineral in the world), which binds potassium very tightly, but not sodium. Sodium, on the other hand, has mostly all been leached out of the
ground and into the oceans eons ago.
The chemistry of feldspar is pretty neat too. Its name means "not ore" or something like that in German, and it mostly contains silica. Aluminum has
a similar atomic radius as silicon, and can fit into a silica matrix, but then it has that missing electron that messes up the matrix. Unless, of
course, potassium sits next to it and lends it its extra electron. Then everything works out great, since potassium fits quite well into that matrix
too. And it means that terrestrial plants have access to an alkali metal as well, because this planet would be a totally different color if there
wasn't one available to them.
Of course, our taste for sodium, combined with the ease of mining it from old dried up seabeds, has meant that we tend to consume it preferentially
over potassium. However, it's not actually bad for us and doesn't contain any calories, so there's no reason to stop, as long as we're getting enough
of all our other minerals too. Sodium nitrite though, is a whole different story. Even though it's food grade, it's a known carcinogen. However,
the FDA has determined that at the very low levels that it's used in meat as a preservative, that the benefits of not getting food poisoning outweigh
the (very low) risk of developing cancer from consuming it.
The first step in the process of learning something is admitting that you don't know it already.
I'm givin' the spam shields max power at full warp, but they just dinna have the power! We're gonna have to evacuate to new forum software!
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AJKOER
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Here is a new approach based on my prior attempts (see http://www.sciencemadness.org/talk/viewthread.php?tid=52&... ) that I will hopefully be able to test soon.
First, as occurs in the case of the metal Aluminum (see, for example, equation (3.7) in a doctoral thesis from 2008, "Alkaline dissolution of
aluminum: surface chemistry and subsurface interfacial phenomena", by Saikat Adhikari, link: https://www.google.com/url?sa=t&source=web&rct=j&...), I would argue similarly with either aluminum or zinc, the creation of the metal
hydroxide directly from the metal, proceeds with the release of electrons per the reaction
Al + 3 OH- → Al(OH)3 + 3 e-
Zn + 2 OH- → Zn(OH)2 + 2 e-
Then, a possible reaction in the presence of nitrate, with either prehydrated or totally solvated electrons, being reported as readily scavenged by
nitrate:
e(p)-/e(aq)- + NO3- + H2O -> NO2 + 2OH- (Source: see eq. (5) in JAERI-Conf 95-003, "5. 6 Radiolysis of Concentrated Nitric Acid Solutions R.
Nagaishi" by P.Y. Jiang, et al, link: https://www.google.com/url?sa=t&source=web&rct=j&... )
Upon shaking the solution periodically, likely containing NO2 gas, in an atmosphere of pure oxygen also (see below):
2 NO2 + H2O --> HNO2 + HNO3
Implying a net reaction of in the case of Aluminum of aqueous nitrate in an alkaline solution:
2 Al + 3 NO3- + 3 H2O -- 6OH- --> 2 Al(OH)3 + 3 NO2-
Do not use more aluminum then needed. But if in excess, expect:
e(p)-/e(aq)- + NO2- + H2O -> NO + 2OH-
and shaking with O2:
2 NO + O2 --> 2 NO2
2 NO2 + H2O --> HNO2 + HNO3
-----------------------------------------------
Possible other reaction of interest would be the formation of the superoxide radical anion (from O2 + e-), which could readily react with any formed
NO, creating peroxynitrate that would be converted back into nitrate.
I also suspect the reaction may proceed well with NH4NO3.
[Edited on 3-10-2017 by AJKOER]
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Σldritch
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| Quote: Originally posted by AJKOER | Here is a new approach based on my prior attempts (see http://www.sciencemadness.org/talk/viewthread.php?tid=52&... ) that I will hopefully be able to test soon.
First, as occurs in the case of the metal Aluminum (see, for example, equation (3.7) in a doctoral thesis from 2008, "Alkaline dissolution of
aluminum: surface chemistry and subsurface interfacial phenomena", by Saikat Adhikari, link: https://www.google.com/url?sa=t&source=web&rct=j&...), I would argue similarly with either aluminum or zinc, the creation of the metal
hydroxide directly from the metal, proceeds with the release of electrons per the reaction
Al + 3 OH- → Al(OH)3 + 3 e-
Zn + 2 OH- → Zn(OH)2 + 2 e-
Then, a possible reaction in the presence of nitrate, with either prehydrated or totally solvated electrons, being reported as readily scavenged by
nitrate:
e(p)-/e(aq)- + NO3- + H2O -> NO2 + 2OH- (Source: see eq. (5) in JAERI-Conf 95-003, "5. 6 Radiolysis of Concentrated Nitric Acid Solutions R.
Nagaishi" by P.Y. Jiang, et al, link: https://www.google.com/url?sa=t&source=web&rct=j&... )
Upon shaking the solution periodically, likely containing NO2 gas, in an atmosphere of pure oxygen also (see below):
2 NO2 + H2O --> HNO2 + HNO3
Implying a net reaction of in the case of Aluminum of aqueous nitrate in an alkaline solution:
2 Al + 3 NO3- + 3 H2O -- 6OH- --> 2 Al(OH)3 + 3 NO2-
Do not use more aluminum then needed. But if in excess, expect:
e(p)-/e(aq)- + NO2- + H2O -> NO + 2OH-
and shaking with O2:
2 NO + O2 --> 2 NO2
2 NO2 + H2O --> HNO2 + HNO3
-----------------------------------------------
Possible other reaction of interest would be the formation of the superoxide radical anion (from O2 + e-), which could readily react with any formed
NO, creating peroxynitrate that would be converted back into nitrate.
I also suspect the reaction may proceed well with NH4NO3.
[Edited on 3-10-2017 by AJKOER] |
I doubt you will get much other than ammonia similar to how you will not get alcohols from clemmensen reduction; the reactant is bonded to the metal
surface until it has picked up hydrogen on all its bonds to the metal so the reaction will not stop at nitrite very often.
If you did it without a solvent it would probably just explode so that is not an option though i think chemplayer did something like this and failed
in a less dramatic way:
https://www.youtube.com/watch?v=w9nhdpKhztI&t=1s
I think lead is about the strongest reducing agent that will work. Other reducing agents that work such as polysulfides and carbon are weaker reducing
agents and both of them react way too fast without something to slow them down. Maybe it depends on the solubility of lead oxide in sodium
nitrate/nitrite too. I doubt aluminium oxide is very soluble but zinc oxide might be.
Then there is this too, i do not know how he/she did the nitrite test but supposedly there is a lot of nitrite in it. NOt sure if if was a commercial
test solution or the ferrous sulfate test though. If it was the ferrous sulfate test than there was not a lot of nitrite it seems.
https://www.youtube.com/watch?v=5Sgd1wjpywc
If you really want to use a metal then i think you should look for one less reducing like bismuth.
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argyrium
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Am I the only one who spotted this error in the original post - or am I in error?
"Then i thought i would neutralize the residual sodium hydroxide so i added sodium bicarbonate repeated the filtering again. I really should have
thought about that earlier but oh well."
??
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