Dionysus12
Harmless
Posts: 3
Registered: 9-7-2011
Member Is Offline
Mood: No Mood
|
|
Preparation of Sodium Nitrite NaNO2
Hello everyone,
I've been meaning to make or procure some sodium nitrite which has suddenly become very hard to find in my country. I was thinking of reacting HCL
with KNO3 and passing the fumes through a bath of aqueous NAOH but I wasn't sure about how to go about seperating the products. Any advice?
|
|
|
barley81
Hazard to Others
Posts: 481
Registered: 9-5-2011
Member Is Offline
Mood: No Mood
|
|
Please use the search engine before asking. There was a thread on the topic. Here it is:
http://www.sciencemadness.org/talk/viewthread.php?tid=52
and here is a method I think is nice:
Quote: Originally posted by Polverone  | The other night I was looking at Muspratt volume 1 under "nitrous ether" (ethyl nitrite) and I found what seems like an interesting, simple method.
Mix 100 parts of very fine KNO3 with 12.07 parts of lampblack (I'm guessing any finely divided carbon would work - I hope so), heat it in a crucible,
and cover it and remove it from the heat source when the reaction is concluded (http://bcis.pacificu.edu/~polverone/muspratt1/c-834.html). According to the book some carbonate is formed but it is mostly pure enough to be used
as-is (at least by the standards of 1860). It also mentions the formation of silicate - not sure if this is because they started with impure KNO3 or
because of the crucible's composition. I am going to try this soon, not because I need more nitrite but because every other preparative method I've
come across has been such a pain.
BTW, kingspaz, what physical form is your aluminum in? Powder, granules, foil, wire?
[Edited on 18-2-2003 by Polverone] |
edit:
I only picked that method because the reactants are cheap and it's a little easier (for a few g of nitrite-containing slag).
[Edited on 9-7-2011 by barley81]
|
|
|
AndersHoveland
Hazard to Other Members, due to repeated speculation and posting of untested highly dangerous procedures!
Posts: 1986
Registered: 2-3-2011
Member Is Offline
Mood: No Mood
|
|
KNO3 is not going to react with HCl, at least not in water (it might be possible to use DMSO or an ionic liquid as the solvent). Assuming you can get
the KNO3 to dissolve, anhydrous HCl gas could be bubbled in and the reaction would probably be:
KNO3 + (4)HCl --> KCl + (2)H2O + NOCl + Cl2
Bubbling the gas products into a solution of NaOH will not give any nitrites, because the chlorine would oxidize them to nitrate.
NaNO2 + H2O + Cl2 --> NaNO3 + (2)HCl [aq]
Possibly it would work if there was also thiosulfate in the solution to immediately reduce the chlorine, but not sure.
The reduction of sodium nitrate to nitrite with carbon is not efficient and gives low yields. Much higher yields of nitrite can be obtained by using
molten lead instead. Alternatively, 25% concentrated nitric acid can react with copper to give off nitrogen-oxide fumes that can then be bubbled to an
alkaline solution to form nitrites. Do not use higher concentration of HNO3, as then mostly nitrogen dioxide will form which will reduce yields.
(2)NaOH + NO2 + NO --> (2)NaNO2 + H2O
Remember that both nitric oxide (NO) and solutions of sodium nitrite (NaNO2) are oxidized by the air, to brown NO2 and NaNO3, respectively.
I'm not saying let's go kill all the stupid people...I'm just saying lets remove all the warning labels and let the problem sort itself out.
|
|
|
Tiberius_C
Harmless
Posts: 9
Registered: 3-11-2010
Member Is Offline
Mood: No Mood
|
|
Other than reacting with carbon, you might want to try warming a small quantity to just past decomposition temp -- should yield NaNO2 
|
|
|
Dionysus12
Harmless
Posts: 3
Registered: 9-7-2011
Member Is Offline
Mood: No Mood
|
|
Thanks for the help, I might try the carbon route after hitting the books for a bit. I tried using the search function but I couldn't find anything
relevant for some reason.
|
|
|
Magpie
lab constructor
Posts: 5939
Registered: 1-11-2003
Location: USA
Member Is Offline
Mood: Chemistry: the subtle science.
|
|
http://www.sciencemadness.org/talk/viewthread.php?tid=52&...
The single most important condition for a successful synthesis is good mixing - Nicodem
|
|
|
Dionysus12
Harmless
Posts: 3
Registered: 9-7-2011
Member Is Offline
Mood: No Mood
|
|
Thanks Magpie, the pictures are very helpful
|
|
|
Chemistry Alchemist
Hazard to Others
Posts: 403
Registered: 2-8-2011
Location: Australia
Member Is Offline
Mood: No Mood
|
|
Heating up Nitrates make them loose a oxygen molecule forming NO2 (2NaNO3 -> 2NaNO2 + O2?)
just a guess, remember reading up on it a year back
|
|
|
redox
Hazard to Others
Posts: 268
Registered: 22-2-2011
Location: The Land of Milk and Honey
Member Is Offline
Mood: Chalcogenetic
|
|
I think that reaction only occurs smoothly in the presence of a reducing agent, such as carbon or lead. If it were as simply as you made it seem,
there wouldn't be huge threads on the topic.
My quite small but growing Youtube Channel: http://www.youtube.com/user/RealChemLabs
Newest video: Synthesis of Chloroform
The difference between chemists and chemical engineers: Chemists use test tubes, chemical engineers use buckets. |
|
|
Chemistry Alchemist
Hazard to Others
Posts: 403
Registered: 2-8-2011
Location: Australia
Member Is Offline
Mood: No Mood
|
|
Could you heat up a test tube or beaker of sodium nitrate, mix it with some Carbon and heat it up? (NaNO3 + C = NaNO2 + CO) is this possible? sorry if
its wrong... im just reading different forums and putting things together
[Edited on 22-9-2011 by Chemistry Alchemist]
|
|
|
ScienceSquirrel
International Hazard
Posts: 1863
Registered: 18-6-2008
Location: Brittany
Member Is Offline
Mood: Dogs are pets but cats are little furry humans with four feet and self determination! 
|
|
I have made potassium nitrite of I believe to be reasonable purity by melting potassium nitrate (mp 334 C ).
It melts to a clear liquid that gradually loses oxygen and solidifies as it forms the nitrite (mp 440 C).
In my opinion this is easier than trying to make the sodium salt as the melting points of sodium nitrate and sodium nitrite are 308 C and 271 C
respectively.
A solution of the product was only weakly basic and addition of a solution of the product to a solution of dilute sulphuric acid and ferrous sulphate
gave a strong brown ring.
[Edited on 22-9-2011 by ScienceSquirrel]
|
|
|
AndersHoveland
Hazard to Other Members, due to repeated speculation and posting of untested highly dangerous procedures!
Posts: 1986
Registered: 2-3-2011
Member Is Offline
Mood: No Mood
|
|
| Quote: Originally posted by ScienceSquirrel | I have made potassium nitrite of I believe to be reasonable purity by melting potassium nitrate (mp 334 C ).
It melts to a clear liquid that gradually loses oxygen and solidifies as it forms the nitrite (mp 440 C).
In my opinion this is easier than trying to make the sodium salt as the melting points of sodium nitrate and sodium nitrite are 308 C and 271 C
respectively.
|
This is very interesting. My only question would be what portion of the potassium nitrate can be converted to nitrite by heating. Potassium nitrate
begins to decompose at a relatively low temperature, but it was my understanding that it is mostly an incomplete decomposition unless much
higher temperatures are reached. That is to say that only a very small portion of the sample being heated will turn into nitrite.
[Edited on 22-9-2011 by AndersHoveland]
|
|
|
Alastair
Hazard to Self
Posts: 59
Registered: 13-7-2011
Member Is Offline
Mood: Barely any solvent in my emulsion
|
|
4189. Nitrite of Potassa. It is obtained
mixed with a little nitre and potash by
heating nitre to redness. To purify the
residuum, dissolve it in boiling water, set
aside for 24 hours, pour off the liquid from
the deposited nitre, neutralize the free alkali
with acetic acid, and add twice its volume of
alcohol. In a few hours more, nitre crystallizes,
and the liquid separates into two layers ;
the upper is alcoholic solution of acetate of
potash, the lower is solution of nitrate of
potash, which may be evaporated to dryness,
or kept in solution. (Bcasley.)
Or, pass nitrous acid gas, formed by acting:
on 1 part of starch with 10 of nitric acid,
through a solution of caustic potash, specific'
gravity 1.38, until it becomes acid ; then add'
a little caustic potash, so as to render it distinctly
alkaline. It may then be kept in the
liquid form, or evaporated to dryness. ( Corenwinder) <-- this works good only in VERY basic conditions.
|
|
|
AJKOER
Radically Dubious
Posts: 3026
Registered: 7-5-2011
Member Is Offline
Mood: No Mood
|
|
Here is an untried but interesting theoretical process to produce small quantities of NaNO2. On a prior Sciencemadness thread:
http://www.sciencemadness.org/talk/viewthread.php?tid=6650
it was noted that NH4NO2 can be formed by the oxidation of a small quantity of NH4OH with H2O2 in the presence of Na2CO3 or NaOH or Zn dust or Pt.
2NH4OH + 3H2O2 --catalyst-> NH4NO2 + 6H2O
So just add NH4OH to say Sodium Percarbonate, 2Na2CO3 . 3H2O2 (found in household chlorine free bleach, but conform to local laws) and bring to a boil
(caution, I would not rule out the possibility of an explosion or toxic fumes upon full evaporation). The presence of the nitrite can be reputedly
demonstrated upon adding H2SO4 to form NO and NO2:
2 NH4NO2 + H2SO4 --> (NH4)2SO4 + 2 HNO2
2 HNO2 --> NO + NO2 + H2O
I suspect that in the presence of Na2CO3 some NaNO2 could be formed:
2 NH4NO2 + Na2CO3 --> (NH4)2CO3 + 2 NaNO2
the reaction moving to the right as the Ammonium carbonate decomposes upon heating at 58 C into NH3, CO2 and H2O (NaNO2 only decomposes above 320 C).
Also, it is likely that some of the NH4NO2 will decompose upon heating in water:
NH4NO2 + H2O --> N2 + 2 H2O
Note, the formation of NaNO2 is generally desirable given the explosive properties of NH4NO2 (both thermal detonation between 60 to 70 C and shock
sensitivity), its 2 hour half-life, and its acutely toxic nature.
[Edited on 3-12-2011 by AJKOER]
[Edited on 3-12-2011 by AJKOER]
|
|
|
Nitro-esteban
Harmless
Posts: 39
Registered: 10-4-2013
Location: Fifth dimension
Member Is Offline
Mood: inert
|
|
I came up with this equation according to some info I got from wikipedia: 2NH3 + 3 H2O2 = NH4NO2 + 4 H2O.
Im thinking that maybe barium nitrite could be made by reacting barium nitrate with ethanol.
|
|
|
subsecret
Hazard to Others
Posts: 424
Registered: 8-6-2013
Location: NW SC, USA
Member Is Offline
Mood: Human Sadness - Julian Casablancas & the Voidz
|
|
I've heard that KNO3 + Cu + HCl will react to form NO2.
Fear is what you get when caution wasn't enough.
|
|
|
SM2
Hazard to Others
Posts: 359
Registered: 8-5-2012
Location: the Irish Springs
Member Is Offline
Mood: Affect
|
|
Just do the litherage (Pb) starting with the NO3+. It consistently works, and the red and yellow lead oxides can be recycled, or used elsewhere.
"Old men who speak of victory
shed light upon their stolen life
they - drive by night- and act as if they're
moved by unheard music." B. Currie
|
|
|
AMchemistry
Harmless
Posts: 3
Registered: 24-7-2013
Location: California
Member Is Offline
Mood: Happy
|
|
you can heat up the NaNO3 to decompose it to NaNO2
-AMchemistry
|
|
|
AndersHoveland
Hazard to Other Members, due to repeated speculation and posting of untested highly dangerous procedures!
Posts: 1986
Registered: 2-3-2011
Member Is Offline
Mood: No Mood
|
|
reduction of nitrates to nitrites using molten lead
see attachments
  
|
|
|
ElectroWin
Hazard to Others
Posts: 224
Registered: 5-3-2011
Member Is Offline
Mood: No Mood
|
|
if you needed food grade nitrite, dont use the lead process
|
|
|
AndersHoveland
Hazard to Other Members, due to repeated speculation and posting of untested highly dangerous procedures!
Posts: 1986
Registered: 2-3-2011
Member Is Offline
Mood: No Mood
|
|
If sodium nitrate is dissolved in concentrated hydrochloric acid, an equilibrium is established:
NO3‒ + 3 Cl‒ + 4 H+(aq) <==> 2 H2O +
NOCl + Cl2
Under strongly acidic conditions, the hydrochloric acid can reduce the nitrate. If such a mixture is heated to boiling and distilled, the fumes could
be passed into a solution of ferrous sulfate, FeSO4, resulting in nitric oxide. That nitric oxide could then be bubbled, along with a
limited amount of air into a third solution containing sodium bicarbonate, to neutralize the mixed nitrogen oxides, resulting in sodium
nitrite.
Assuming the solution is not too acidic, nitrosyl chloride hydrolyzes into dilute hydrochloric and nitrous acids:
NOCl + H2O --> HNO2 + H+(aq) + Cl‒
2 Fe2+(aq) + Cl2 --> 2 Fe3+(aq) + 2 Cl‒
Iron(II) salts will not reduce nitric oxide, but they will reduce nitrogen dioxide to nitric oxide.
Nitric oxide oxidizes on contact with air to nitrogen dioxide. If an equal volume mix of these gases (1 part NO2, 1 part NO) are passed into an
alkaline solution, nitrite will form.
2 NaOH (aq) + NO2 + NO --> 2 NaNO2 + H2O
Of course, for maximum efficiency (and purity), we do not want all of the NO converted into NO2. Let us look at what would happen if too much NO2 was
allowed to react...
2 NaOH (aq) + 2 NO2 --> NaNO3 + NaNO2 + H2O
If the proportion of NO2 is too great relative to the amount of base in solution, none of the NO will be absorbed, and no nitrite will be able to
form.
NaNO2 + NO2 --> NaNO3 + NO
Furthermore, your newly made solution of sodium nitrite needs to be kept away from air, because it will oxidize.
So, it is important to make sure that too much air is not allowed to react, and it may be best to use a slight excess of base. Sodium nitrite
crystallizes as white cubic crystals, so it is fairly easy to identify.
An alternative way, without using air, would be to use dilute nitric acid to convert some of the nitric oxide into nitrogen dioxide. This would
require a fourth solution in the setup. Get out all your stoppers and tubing! (Did I mention that NO2 can be a bit corrosive to rubber
hoses?)
2 NO3‒ + 2 H+(aq) + NO --> H2O + 3 NO2
Normally the nitrate ion is fairly inert to reduction, but nitric oxide is a reactive radical that can catalyze many reactions. A little acid
is required, however. If conditions were alkaline, just the reverse reaction would occur, NO2 can oxidize nitrite to nitrate.
Solutions of nitrous acid exist in equilibrium with nitric oxide and nitrogen dioxide. Furthermore, nitrogen dioxide is considered a mixed acid
anhydride because in water it reversibly hydrolyzes into nitrous and nitric acids.
I am not sure, you may just be able to get the concentrated hydrochloric acid, nitrate salt, and iron(II) salt to all react
together to directly give off the nitric oxide. And then carefully bubble this into the bicarbonate solution along with a small flow of air, being
very careful only to let in a very small amount of air. Combined, the two equations would look like this:
NaNO3 + 3 HCl + 3 FeCl2 --> NaCl + 3 FeCl3 + NO
4 NO + O2 + 2 NaCO3H --> 2 NaNO2 + H2O + 2 CO2
[Edited on 28-7-2013 by AndersHoveland]
|
|
|
AJKOER
Radically Dubious
Posts: 3026
Registered: 7-5-2011
Member Is Offline
Mood: No Mood
|
|
To confirm Andershoveland's last equation,
4 NO + O2 + 2 NaHCO3 --> 2 NaNO2 + H2O + 2 CO2
Atomistry.com commenting on the preparation of Sodium nitrite lists one of the paths as:
"; the action of nitrous fumes containing excess of nitric oxide on a solution of sodium hydroxide or carbonate"
Link: http://sodium.atomistry.com/sodium_nitrite.html
---------------------------------------------------------------------------------------
The same source also notes:
"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, whereas
the action of coal and graphite is too energetic:
2NaNO3 + Ca(OH)2 + C= 2NaNO2 + CaCO3 + H2O "
For those new to the literature, the comment 'too energetic' above could mean it acts like an explosive (detonates, see http://www.sciencemadness.org/talk/viewthread.php?tid=21101&... ) especially in a closed container, so be wary if attempting such a path in any
significant quantities.
[Edited on 31-7-2013 by AJKOER]
|
|
|
symboom
International Hazard
Posts: 1143
Registered: 11-11-2010
Location: Wrongplanet
Member Is Offline
Mood: Doing science while it is still legal since 2010
|
|
Sodium nitrite
Sodium nitrite from ammonia on sodium percarbonate
Ammonia and peroxide form ammonium nitrite
Ammonium nitrite and sodium carbonate
With heating decomposes ammonium carbonate from solution
Into ammonia and co2 gas
[Edited on 7-10-2016 by symboom]
[Edited on 7-10-2016 by symboom]
|
|
|
Texium
|
Thread Moved
7-10-2016 at 08:50 |
clearly_not_atara
International Hazard
Posts: 2787
Registered: 3-11-2013
Member Is Offline
Mood: Big
|
|
I like your idea, AndersHoveland, but I think you could avoid gases more, which is always nice when equipment is limited. Maybe you could condense
most of the NOCl with some of the Cl2 in cold dichloromethane or chloroform (NOCl should have no chance of oxidizing chloroform without activation).
Then passing it over eg copper, ferrous sulfate, or possibly eg oleic acid, anything that reacts heterogeneously with Cl2 could either purify the NOCl
while reducing Cl2 or release NO gas.
Really I think finding something that will react with Cl2 but not NOCl is the trick. Maybe CuCl? NOCl in CH2Cl2 is a useful end-product
|
|
|
![[*]](images/xpblue/default_icon.gif){kind=icon}
