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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.
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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.
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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)
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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...
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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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Lionel Spanner
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Having made 7 attempts at it, 5 in tin cans and 2 in a stainless steel saucepan, I can only conclude the Grossmann method for making inorganic
nitrites (US Patent 792,515, 1905) which is cited in the wiki, is not useful or reliable. The results vary wildly depending on the surface material of
the reaction vessel, and when successful (in tin cans previously used for food) the nitrite was of very low purity, i.e. 40% or less. The two runs in
stainless steel produced no nitrite at all; the only product recovered was unreacted sodium nitrate.
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Myc
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I have been pursuing the old molten lead method over the last couple of weeks. 85 g Sodium nitrate and 207 g lead (I used lead fishing sinkers) were
melted in a cast iron pot and stirred continuously. An orange oxide begins forming, and the mixture bubbles. After about 30 minutes, the mixture is a
thick orange mud and the lead seems to have mostly reacted. Stirring is continued after the heat is stopped to prevent a solid mass forming. Once the
mixture solidifies, about 200ml water is added and the mixture left to soak for 15 minutes or so. All chunks should disintegrate.
The mixture is filtered and CO2 is bubbled through the filtrate for a few minutes; a white precipitate forms, which is removed by filtration. The
filtrate is then reduced by boiling. When the temperature reaches 125-130 degrees C, heat is removed and the mixture cooled to fridge temperature. The
crystals are then filtered and the filtrate reduced again for a second crop. Adding HCl to the salt gives plumes of brown gas. Pretty straightforward,
right?
Now it was clear to me from early on that not all the lead had reacted. There were small grains of metallic lead amongst the lead oxide. So I tested
the purity of my product as follows. 1 g of my salt was dissolved in a couple of ml of water. 2.5 g silver nitrate was similarly dissolved separately
in 5 ml or so. The two solutions were mixed; a white precipitate of silver nitrite formed instantaneously. This was filtered, dried and weighed. If
the 1g was pure sodium nitrite, the yield of silver nitrite should be 2.23g.
After one lead/nitrate reduction, my salt mixture was 30% nitrite. I then repeated the procedure with this salt mixture and fresh lead and got it to
50%. A third cycle got me to 80%. For reference I also tested some sodium nitrite isolated from curing salt, which I measured at 95% (although, in
this case, I think any sodium chloride contaminant would have formed the insoluble silver chloride which could throw my numbers off).
It's possible that a longer reaction could boost yield, however I'm also aware that sodium nitrite decomposes into the nitrate in the presence of
oxygen at high temperatures, so a longer reaction time could possibly be counterproductive. So all in all, to me, the lead method is 'not great'.
While I haven't tried it, the use of silver nitrate could also be used as a purification method if you're happy to buy/make plenty of silver nitrate
(you'll need 250g of silver nitrate to separate 100g of the sodium nitrite from nitrate contamination, but you can reclaim most of it!). After
reacting as above, the silver nitrite would be mixed with an equimolar amount of sodium chloride in solution to give a precipitate of silver chloride,
while sodium nitrite remains in solution. Filter the solid and evaporate the water for your sodium nitrite. The silver chloride can be made back into
metallic silver with NaOH and sugar, then reacted with nitric acid to regenerate the silver nitrate.
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Lionel Spanner
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For what it's worth I've recently tried a couple of methods described in this thread, with no joy.
The reaction of sulphamic acid, calcium oxide and sodium nitrate (as described on the first page) resulted in a lot of water vapour, some nitrogen dioxide, and the precipitation of a non-reducing substance that isn't nitrite
or nitrate, and is considerably less soluble in water than simple inorganic nitrites.
The reaction of sodium nitrate and sodium sulphide (as described in Morgan, 1908) produced sodium sulphate, and a hygyroscopic yellow product that was
slightly less soluble in water than nitrate or nitrite, and was contaminated with unreacted sodium sulphide - possibly sodium sulphamate.
Given that inorganic homebrewed nitrite is the modern-day amateur chemists' equivalent of the philosopher's stone, I'm very glad that Poland exists,
Polish vendors will freely sell sodium nitrite to private individuals, and although it's not cheap, the postage is not ridiculously expensive.
[Edited on 19-2-2023 by Lionel Spanner]
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Lionel Spanner
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Many years ago I used to work at a cosmetic/toiletry manufacturer that had once used bronopol (2-bromo-2-nitro-propane-1,3-diol) to preserve many of
its products, and found some of them turned brown to black due to the reaction between nitrite ions released by bronopol degradation and cocamide DEA,
turning the latter into an unstable N-nitrosamine due to an incomplete diazotisation reaction; when attempted with secondary amines, this reaction
stops at the nitrosamine intermediate. (As nitrosamines are highly carcinogenic, this was extremely bad news, and the preservative system in those
products was soon changed.)
As it turns out, bronopol is much more rapidly hydrolysed to nitrite in aqueous caustic soda at 100 °C. The initial products of the reaction,
formaldehyde and 2-bromo-2-nitroethanol, are relatively volatile, boiling at -19 and 83 °C. This could potentially be turned into a useful
preparation, though bronopol is hard to come by for amateurs.
Source: Sanyal, Basu, Banerjee. Rapid ultraviolet spectrophotometric determination of bronopol: application to raw material analysis and kinetic
studies of bronopol degradation. , J. Pharm. Biomed Anal., 14 (1996), 1447–1453. doi:10.1016/0731-7085(96)01779-7
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