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AJKOER
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Reagent Friendly Path: Just a small starting amount of H2O2, aqueous NH3, Cu, air and some NaOH to maintain alkaline pH, and as this is an
electrochemical cell, a small touch of a good stating electrolyte like sea salt, all resulting in NaNO2.
It is well known that the electrochemical dissolving of Cu with aqueous NH3 in the presence of air creates a side product of NH4NO2, which apparently,
at one point (I suspect removal of ammonia resulting in a lower pH), creates a massive sudden N2 release per the decomposition of formed NH4NO2 (so,
do employ a wide mouth and tall reaction vessel and still, at times, I get a spillage event).
See my prior comments on this event and details on the reaction previously posted here https://www.sciencemadness.org/whisper/viewthread.php?tid=18... which also includes another reference on the electrolysis route to nitrite (not a
galvanic cell).
Now, adding NaOH likely averts this NH4NO2 decomposition reaction forming the targeted NaNO2.
Also, the reaction with ammonia is depicted as proceeding both with H2O2 and O2 exposure along I suspect, it is not O2, but perhaps superoxide formed
from solvated electrons (e- per the electrochemically cell) acting on oxygen (O2 + e- -> .O2-). Basis, quoting from a source at https://www.frontiersin.org/articles/10.3389/fmars.2016.0023... :
"Superoxide (O−2) also reacts with trace metals (Figure 1B) to produce H2O2; reactions with Cu(I) can proceed rapidly with rate constants (k) on the
order of 2 × 109 M−1 s−1 (Zafiriou et al., 1998), although [Cu(I)] is relatively low (~0.1 nM; Moffett and Zika, 1988)."
which is my polite way of saying that a direct oxygen path may not be precisely correct.
[Edited on 16-7-2022 by AJKOER]
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Fantasma4500
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AJKROER
you mention formation of nitrite in the same reaction vessel as.. superoxide, and also using H2O2? nitrite reacts with oxygen to form nitrate
if this is a membrane cell it adds difficulty to the operation, though its very possible to make your own membrances with PVC glue and plastic cloth
would be cool if it was a viable electrochemical route for making nitrite
i would go with oswald if i had space to have such a thing going, simply dumping the gasses through water and then have them pumped through hydroxide
solution should do
the starch method is interesting, any mentions of what may be formed? nitrostarch is a thing..
im not quite sure of how one would gauge the quality of the gas mixture, amount of gas bubbles in second container after first water-scrub? im not
even sure if NOx from nitric acid is viable for this.
i employed my IPN method, i got 100mL of presumably mostly IPN
so i set up for short path distillation and ended up with a majestic.... 35 milliliters of IPN, from maybe 300mL 62% HNO3 - hmmmm, im not sure if this
is a good yield.
i did insert an airbubbler into my flask to counteract the ups and downs in temperature from my hotplate and the reaction- afaik nitrite isnt super
sensitive to oxygen, but maybe im wrong, maybe i turned my poor nitrite into nitrate like that. maybe i messed up on making the IPN itself?
regardless if one goes with the NOx method, its ideal to pump the gasses into KOH solution, then dumping it in freezer will crystallize the KNO3 and
youre left with mainly KNO2 - K2CO3 may also be used
how would one go about modifying the starch reaction, is the nitric acid dripped into solid starch, into starch solution, starch solution dripped in?
controlled heating?
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AJKOER
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Correctly, there are claims of the direct action of oxygen gas on nitrite (and not superoxide) with NaNO2, as an example, a decomposition warning for
the dry salt (see https://nj.gov/health/eoh/rtkweb/documents/fs/2258.pdf ), to quote:
"away from AIR, LIGHT, and MOISTURE."
However, I suspect, this trilogy are hardly independent agents, as nitrites are noted as photocatalysts.
NO2- + Light (Blue) -> .NO2 + e-
And, the hydrolysis of .NO2 is a known path to nitrous and nitric acid:
.NO2 + .NO2 + H2O -> HNO2 + HNO3
Again the action of light on nitrous acid:
HONO + Light -> .HO + .NO
And, the action of a solvated electron on oxygen:
O2 + e-(aq) --> .O2- (the introduction of the superoxide radical anion in the presence of an oxygen source)
which can also be sourced in the present of a transition metal impurity, in the so-called metal auto-oxidation reaction with dissolved oxygen:
M -> M+ + e-
O2 (aq)+ e- --> .O2- (aq) (again with the formation of superoxide)
And, the action of superoxide on nitric oxide is a well known path to peroxynitrite and eventually nitrate:
.NO + .O2- -> ONOO
So, once the photo or a transition metal inducing electron presence ceases (not just air presence, for example), a nitrite product could be stable.
This is supported here in the copper, ammonia, O2/H2O2 reaction system with no citations of any ammonium nitrate creation.
Also, there is a source citing the presence of ammonium carbonate (see https://www.sciencemadness.org/whisper/viewthread.php?tid=14... ) fostering the dissolving of copper ore with water, ammonia and oxygen. This is
interesting as CO2/HCO3- appears to be a radical reaction promoter, to quote a source ("Radical production by hydrogen peroxide/bicarbonate and copper
uptake in mammalian cells: Modulation by Cu(II) complexes" at https://www.sciencedirect.com/science/article/pii/S016201341... ):
"It is well-known that the bicarbonate/carbon dioxide pair, the presence of which is important in maintaining physiological pH in extracellular body
fluids, can accelerate the transition metal ion-catalysed oxidation of various biotargets. Despite of its relevance, however, most of the mechanisms
that have been proposed to account for this important effect remain controversial [8], [9], [10], [11], [12], [13], [14], [15], [16], [17], [18],
[19], [20], [21]. On the other hand, it is accepted that the bicarbonate/carbon dioxide pair can increase peroxynitrite-mediated one-electron
oxidation and nitration via formation of the carbonate radical and nitrogen dioxide [22], [23].
So discounting possible underlying radical based reaction paths (aka, not just simply a direct elemental oxygen interaction) is not likely precisely
correct, in my opinion.
As further evidence of the actual possible complexity in systems involving dissolved (not gaseous) .NO see my sources cited relatedly in this SM
thread https://www.sciencemadness.org/whisper/viewthread.php?tid=15... where dissolved .NO apparently behaves differently from gaseous .NO interacting
with oxygen with a near total absence of eventual nitrate creation.
[Edited on 23-7-2022 by AJKOER]
[Edited on 23-7-2022 by AJKOER]
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Lionel Spanner
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Your man Experimental Chemistry has a nice video on preparing potassium nitrite via calcium formate.
https://www.youtube.com/watch?v=BhBxDkU7AmI
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Fantasma4500
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wow. 68% yield of KNO2 starting from Calcium Formate and KNO3, heated at about 300*C, no explosively exothermic reaction
sodium formate may be acquired in 25kg bags as a special de-icer, its used especially for areas that gets very cold but also to avoid salt getting
onto vehicles, aircrafts etc.
sodium nitrate should also be very much doable
2KNO3 + (HCOO)2Ca = 2KNO2 + CaCO3 + H2O + CO2
this seems to be exactly what we have been hoping for to pop up in regards to reduction of nitrate salt. could formic acid maybe reduce nitric acid to
nitrous acid?
The Reaction of nitric acid with formaldehyde and with formic acid and its application to the removal of nitric acid from mixtures
https://onlinelibrary.wiley.com/doi/abs/10.1002/jctb.5010080...
"Nitrites do not react with formaldehyde in neutral solution, " HNO2 may infact be formed in this reaction, it seems.
i see a comment in the calcium formate method video
"nitrite into HCL gives largely nitrosyl chloride"
would this not imply that it can maybe go the other way around again, so NOCl + NaOH = NaNO2 + NaCl?
HCl + KNO3 = NOCl (basically)
anyhow back to calcium formate method:
calcium formate is about 16g/100mL (roughly same 0-100*C)
sodium formate is 49-160g/100mL
CaCl2 is 60 to 160g/100mL
CaCl2 + NaForm = CaForm + NaCl
fractional crystallization would work- not that NaCl would really be a big issue as impurity by my assumptions
could we maybe directly use sodium formate for this reduction instead?
thank you very much for this input @Lionel Spanner
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Fantasma4500
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Ca(NO3)2 + 2NaForm = CaForm + 2NaNO3
164g + 70x 2 140g = BOIL THIS DRY, combust
before / after total weight of dry substance
304g = 195g
i react it at about 360*C, it takes maybe 10 minutes before it starts to react, it doesnt evolve a lot of smoke but it does smell like a nitrate
pyrotechnic composition so thats very annoying if you react too much at once.
https://gyazo.com/c414de838c4d0d7ac40e7399deb23d6f
(COOH)2Ca + 2 NaNO3 = 2 NaNO2 + CaCO3 + CO2 + H2O
130g + 170g = 138g + 100g
dissolve in water, boil NaNO2 dry or use as solution
i have tested a smaller sample of 10 grammes with IPA and HCl
when the IPN is formed it makes the polarity of the IPA seperate out so you get a very clear indication
https://gyazo.com/95df06a3975a0ba352fecd3a3db5bebb
thereafter i ignited the gasses in the flask and typical nitrite flame was seen. its also vasodilating.
its ideal to dissolve the soluble contents in water before adding acid as it causes a lot of effervescence, namely HCl
it may be possible to take a concentrated NaNO2 solution and dump into EtOH to precipitate out the NaNO2 for easy isolation as NaNO2 is 4.4g/100mL
solubility in EtOH
this procedure can be done inside if done in small quantities, 100 grammes was too much
my heating device is a single hotplate, a stainless steel pot ontop of that which is isolated with Al2O3 ceramic wool, secured by a bolt + washer +
nut going through top of the pot
i shall attempt further to simply react Ca(NO3)2 and NaFormate by direct decomposition so i dont have to dissolve that in water, and then boil that
dry
i may add some Al2O3 because the fertilizer i get my Ca(NO3)2 is a mix of ammonium salt and Ca(NO3)2 roughly 90-10% ammonium salt being the latter,
which can decompose... very rapidly on an unfortunate day- i might just do this reaction with homemade Ca(NO3)2 just to be completely sure, energetics
can be very powerful especially if situated on a hotplate covered by a pot
i have no guesstimates for yields yet but this appears to be how were gonna be making nitrites in the future
only things i have to add is that calcium formate decomposes thermally into calcium oxalate- so calcium oxalate should maybe be attempted with sodium
nitrate?
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Myc
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I tried a small batch of the above method (sodium formate and calcium nitrate) too. The reaction in the crucible was rather enthusiastic and measured
over 320C. While some of the nitrite salt survived this I suspect there was also significant decomposition at this temperature (this supposedly
happens above about 300C). So I guess the product has sodium oxide/hydroxide in there too?
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Σldritch
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Hm, is oxalic acid available dirt cheap as some product? I've haven't had much luck, but, I've considered making sulfuric acid from Ferrous Sulfate
(which is dirt cheap itself) with it. The by-product would be Ferrous Oxalate Dihydrate. Maybe worth trying the reaction with that as well if one
could end up with a lot of it as a byproduct...
EDIT: Huh, apparently Calcium Oxalate is almost three orders of magnitude less soluble than Calcium Sulfate.
[Edited on 25-9-2022 by Σldritch]
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B(a)P
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In Australia it is sold in hardware stores as rust and stain cleaner and costs about $15/kg.
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clearly_not_atara
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I was able to order oxalic acid per se, no questions asked, for the cement experiment I still haven't finished. It's sitting in a storage
locker 1000 miles away right now for reasons. It's very useful for cleaning stuff; I think it was labeled as mold prevention for decks or something.
But I wouldn't use it for this because calcium oxalate is pretty inert and won't dissolve in anything. Plus, oxalic acid is strong enough (100 times
stronger than formic acid) to protonate nitrate to some extent and may evolve some NO2.
[Edited on 04-20-1969 by clearly_not_atara]
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Myc
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Hi team,
It seems to me that we have a few different methods in this thread that yield a mixture of sodium nitrate and nitrite. However, there has been little
discussion of separating the two. They seem to have similar solubilities in water. Crystallization by cooling a saturated aqueous solution would thus
work to concentrate the salt that's present in a significantly larger quantity.
Does anyone have ideas for separating a mixture that has a more even ratio of nitrate/nitrite?
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Sir_Gawain
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I've thought of a way to make nearly pure NaNO2. First, you produce crude NaNO2 using one of the reduction methods (like the sulfur/NaOH/NaNO3
reaction). Then dissolve the impure product in water and drip it into a concentrated acid. The gasses (NO and NO2) are led into a cold sodium
hydroxide solution. The reactions are:
2NaNO2 + H2SO4 = Na2HSO4 + NO + NO2 + H2O
Then; NO + NO2 + 2NaOH = 2NaNO2 + H2O
When the sodium hydroxide (now nitrite) solution pH reaches neutral, stop the gas production and evaporate the water to recover your pure product.
“Alchemy is trying to turn things yellow; chemistry is trying to avoid things turning yellow.” -Tom deP.
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Myc
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I believe I've read that this also produces nitrate...
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Sir_Gawain
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Bubbling just NO2 into sodium hydroxide produces both NaNO2 and NaNO3. (2NO2 + 2NaOH = NaNO2 + NaNO3 + H2O)
Using a mixture of NO2 and NO results in only NaNO2.
“Alchemy is trying to turn things yellow; chemistry is trying to avoid things turning yellow.” -Tom deP.
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clearly_not_atara
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Adding NiCl2 to a solution containing KNO2 and KNO3 should precipitate K4Ni(NO2)6*H2O selectively as a brown powder. It may be possible to return this
to KNO2 by reaction with KOH.
[Edited on 04-20-1969 by clearly_not_atara]
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Fantasma4500
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i just have some more input for this.
aluminium may be utilized for this reaction- very dilute HNO3 reacts with Al, somehow bypasses passivation layer- this could maybe imply NO formation
over NO2? think of set&forget kinda reaction.
i got around this as i remembered Al reacts with NO2 to form- what? i didnt get to that, possibly it forms Al(NO2)2 but the Al2O3 which forms from the
HNO3 might bump that into Al(NO3)2 - hm. more concentrated acid may be used with a small amount of HCl being added in- or H2SO4 maybe? HAc?
Phosphoric?
this is a potential low cost method, otherwise birkeland eyde would be better, and yet better would be formate with nitrate
@sir_gawain
why would a MIXTURE of NO2 and NO produce pure NaNO2?
NaOH + NO = NaNO2
NaOH + NO2 = NaNO2+NaNO3
the NO doesnt act reducing
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Sir_Gawain
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NaOH + NO does not make NaNO2 (unless by NaOH + NO = NaNO2 + H). And in this case, NO almost does act as a reducing agent. My guess is that NO2 and
NaOH react to form nitrate and nitrite, then NaNO3 + 2NaOH + 2NO = 3NaNO2 + H2O.
“Alchemy is trying to turn things yellow; chemistry is trying to avoid things turning yellow.” -Tom deP.
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clearly_not_atara
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You could see a reaction like
4 NO + 2 OH- >> N2O + 2 NO2- + H2O
As to whether this actually occurs, I don't know, but Wikipedia seems to think it does. One possible reaction pathway is:
NO* + OH- >> -ONOH
-ONOH + NO >> NO2- + HNO
2 HNO >> H2O + N2O (known reaction)
[Edited on 04-20-1969 by clearly_not_atara]
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Lionel Spanner
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Anyone else had grief trying to boil sodium nitrite solutions to dryness?
I've been able to concentrate it, and get a viscous, highly concentrated solution full of solids with a boiling point around 158 °C - the very last
stage, which involves getting the last bit of water out by drying it an oven at 200 °C is where it's gone tits up, every single time.
1st attempt - used a porcelain dish; the salt crust was so hard I ended up cracking the dish and contaminating the product with bits of porcelain.
2nd attempt - used a dish lined with silicone coated oven-safe aluminium foil; after the concentrated solution was added, the coating lasted a whole 5
seconds before both the product and the foil committed suicide.
3rd attempt - used a small steel bowl; product reacted with the iron in the steel, causing it to evolve nitrogen and foam profusely, spilling much of
it onto the base of the oven, and the final product was heavily contaminated with iron oxides.
Maybe something like an oven-safe silicone mould is the answer....
[Edited on 21-11-2022 by Lionel Spanner]
[Edited on 21-11-2022 by Lionel Spanner]
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Myc
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If you're boiling it down to concentrate it, stop at around 130 degrees then cool it to fridge temperature. The sodium nitrite will crystallise and
you'll have a fairly small amount of water. Vacuum filter this paste while still cold. The filtrate won't have too much product in it, but you could
boil and concentrate again if you have a fair bit. The solid can then be dried under a fan if you have low humidity. Finish the drying process in a
desiccator (sealed box with plenty of dry CaCl).
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Lionel Spanner
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Quote: Originally posted by Myc  | | If you're boiling it down to concentrate it, stop at around 130 degrees then cool it to fridge temperature. The sodium nitrite will crystallise and
you'll have a fairly small amount of water. Vacuum filter this paste while still cold. The filtrate won't have too much product in it, but you could
boil and concentrate again if you have a fair bit. The solid can then be dried under a fan if you have low humidity. Finish the drying process in a
desiccator (sealed box with plenty of dry CaCl). |
Thank you! Since posting my mini-rant I've had time to consider this problem in a calmer manner, and this proposition has a lot in common with the
course of action I'm considering.
Solutions of sodium nitrite become supersaturated on boiling when their boiling point reaches around 126-130 °C, so today I made a fresh batch and
transferred the concentrated distillate to another container once the boiling point got to around 130 °C, and cooled it in the fridge. Net result: a
nicely mobile salt crust in a little bit of pale yellow water.
However, instead of vacuum filtration and dessication, I'm intending to remove the remaining water by azeotropic distillation with xylene (the
azeotrope boiling at 92 °C, and comprised of 40% water/60% xylene at atmospheric pressure.) If this works, then the freshly precipitated salt could
nucleate upon the surface of the existing solids, and the final product, being completely insoluble in xylene, could be separated by filtration and
dried by heat. Fingers crossed!
[Edited on 23-11-2022 by Lionel Spanner]
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clearly_not_atara
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I would cross my fingers too when heating a flammable solvent with a somewhat unstable oxidant...
[Edited on 04-20-1969 by clearly_not_atara]
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Lionel Spanner
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Update: azeotropic distillation with xylene worked up to a point, removing all but about 5% of the water, but did not break up the concentrate as
hoped, most likely due to the differences in density and polarity between the two.
The distillation also became less and less effective as the water content decreased. I managed to remove nearly all of the wet nitrite from the flask
while it was liquid, and place it in a desiccator (read: tightly sealed Tupperware-style box) lined with solid caustic soda.
After 6 weeks in storage, no further weight loss was observed, and the final result was 22.7 g sodium nitrite of unknown purity, having started from
42.5 g sodium nitrate, representing a maximum yield of 66%.
Since dehydrating the product is such a chore, I will attempt to recover it by recrystallisation on the next attempt, most likely from alcohol with
some added water.
[Edited on 4-1-2023 by Lionel Spanner]
[Edited on 4-1-2023 by Lionel Spanner]
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Fantasma4500
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@Lionel Spanner
NaNO2 EtOH 4.4g/100mL, i bet acetone is much worse at dissolving it
precipitate it out by adding in a solvent, then vacuum filter
to speed things up you may blow hot air onto the mouth of the vacuum filter to remove the acetone faster
ideally you would place a beaker with the damp NaNO2 crystals in an oven, sealed up so air doesnt get into it- maybe hotplate. oven set above the
boiling point of the solvent used
air can turn NaNO2 into NaNO3
i must also stress that using a solvent to flush with causes the water to be removed, and heating will now not so much cause the material to dissolve
and gradually form a hard lump, this is especially important if you desire to dump it into maybe an erlenmeyer flask, youre skipping the part where
you have to crush up hard lumps, and with luck you get fine powder right away.
anyhow, NO formation can be done with ammonia and oxygen, air pump into NH4OH solution, then this pumped through Cr2O3, maybe a glass tube of kitty
litter silica gel with Cr2O3- then this lead into NaOH
https://edu.rsc.org/exhibition-chemistry/chromiumiii-oxide-c...
@clearly_not_atara
ive flushed chlorates with acetone many times, a method for making pure NaClO4 is to use acetone as solvent
that solvent has to leave the NaClO4 eventually somehow. NaNO2 is barely an oxidizer, its so weak oxidizer its also a reducing agent
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Lionel Spanner
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| Quote: Originally posted by Antiswat | @Lionel Spanner
NaNO2 EtOH 4.4g/100mL, i bet acetone is much worse at dissolving it
precipitate it out by adding in a solvent, then vacuum filter
to speed things up you may blow hot air onto the mouth of the vacuum filter to remove the acetone faster
ideally you would place a beaker with the damp NaNO2 crystals in an oven, sealed up so air doesnt get into it- maybe hotplate. oven set above the
boiling point of the solvent used
air can turn NaNO2 into NaNO3
i must also stress that using a solvent to flush with causes the water to be removed, and heating will now not so much cause the material to dissolve
and gradually form a hard lump, this is especially important if you desire to dump it into maybe an erlenmeyer flask, youre skipping the part where
you have to crush up hard lumps, and with luck you get fine powder right away. |
I will definitely avoid unnecessary exposure of nitrite to air at high temperatures. It might well be worth removing the bulk of the water by vacuum
distillation, in lieu of a rotary evaporator.
The solubility figure for alcohol is at 25 °C, and it's almost certainly higher at 70-80 °C; acetone could work very well as an anti-solvent, in
order to get as much of it out of solution as possible. The catch is that any unreacted nitrate has a very similar solubility profile to nitrite.
Apparently, nitrite is a little less soluble in water than nitrate at high temperatures (160 vs 180 grams per 100 g water), so if push comes to shove,
the two could be separated by fractional crystallisation, though I'm hoping it won't come to that.
Coming from a UK state school that spent its entire budget on languages, I never really understood the process of recrystallisation, and so I did a
terrible job of it when I was at university. Now I can do it at my leisure, I now understand the process much better, and it just seems so much more
simple.
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