| Pages:
1
..
7
8
9
10
11
..
16 |
woelen
Super Administrator
Posts: 7976
Registered: 20-8-2005
Location: Netherlands
Member Is Offline
Mood: interested
|
|
Trying to use iron in aqueous solution will not produce any nitrite. Nitrate ion is remarkably inert in aqueous solution, except in strongly
concentrated acidic environments or at very high pH in combination with magnesium and aluminium. In the acidic environments it forms NOx with suitable
reductors, in the strongly alkaline environment it forms ammonia.
I have no experience with mixing iron and KNO3 and trying to react that. Just try it on a small scale and report back on your results.
|
|
|
Polverone
Now celebrating 21 years of madness
|
Threads Merged
16-4-2012 at 22:26 |
dann2
International Hazard
Posts: 1523
Registered: 31-1-2007
Member Is Offline
Mood: No Mood
|
|
From another thread.
Making Nitrite from Nitrate + Lead
Attachment: nitrite.zip (213kB)
This file has been downloaded 838 times
|
|
|
chochu3
Hazard to Others
Posts: 185
Registered: 21-10-2003
Location: South Side Tejas [Cloverland]
Member Is Offline
Mood: Inside looking out
|
|
If ionic solutions are preferred some older text which I do not remember have stated it can be done with hydriodic acid and some other substrates, if
I can find my journal with references I have other written formulas which I will post later.
\"Abiding in the midst of ignorance, thinking themselves wise and learned, fools go aimlessly hither and thither, like blind led by the blind.\" -
Katha Upanishad
|
|
|
AJKOER
Radically Dubious
Posts: 3026
Registered: 7-5-2011
Member Is Offline
Mood: No Mood
|
|
Expanding on a previous reference to a 2009 Sciencemadness thread containing a reply by Madcedar entitled "NaNO2 from Ethanol, HNO3 and NaOH"
(https://www.sciencemadness.org/whisper/viewthread.php?tid=52... ), here is an old work dedicated to Ethyl nitrite. Reacting with KOH in alcohol,
for example, yields KNO2.
Source: "Spirit of nitrous ether and ethyl nitrite: a bibliographical index of ...", by W. O. Richtmann, J. A. Anderson. Some extracts:
Link:
http://books.google.com.bz/books?pg=PA34&id=7yfrAAAAMAAJ...
[SEE PAGE 36]
"Persoz prepares ethyl nitrite by the action of pure fuming nitric acid upon absolute alcohol, well cooled, to prevent explosion. The ethyl nitrite is
removed, and purified in the usual manner. 1862."
"The ethyl nitrite is decomposed by caustic potash (especially in alcoholic solution) and yields potassium nitrite and alcohol, proving that ethyl
nitrite is a true "ether." Kopp, E. 1851."
Revue scientif., 27, p. 273. [Jour. de Pharm., 11, p. 320; Ann. d. Chem., 77, p. 332; Neues Jour. de Pharm, Jl, p. 320; Gmelin, vol. 8, p. 470;
Jahresber. u. d. Ftschr. d. Pharmacog.,
2, p. 325; Am. Jour. Pharm., 59, p. 484.] Salpetrigsaures Aethyloxyd.
"Kopp prepares ethyl nitrite by using a mixture of alcohol, nitric acid, copper turnings, and passing the gas through a solution of ferrous sulphate
to remove nitrous acid and other oxides of nitrogen. Gerhardt. 1851."
[SEE PAGE 35]
"Gmelin, vol. 8, p. 468. Nitrite of Ethyl.
History, synonomy, methods of formation, preparation—
1. Fuming nitric acid and alcohol, cold.
2. Fuming nitric acid, alcohol and water.
3. Nitric acid and alcohol, distilling.
4. Sulphuric acid, alcohol and nitrate.
5. Nitric acid, alcohol and reducing substances.
6. Nitrous acid on alcohol direct.
7. Sulphuric acid, alcohol and saltpeter (ord. temp.), purification, properties, decomposition and combinations.
Lea, M. C. 1861.
Sillim. Am. Journ. Sc, 32, p. 177. [Chem. Centrbl., 33, p. 688; Am. Journ. Ph., 34, p. 69; Proc A. Ph. A., 10, p. 160; Ann. der Pharm., 165, p. 58;
Drug. Circ, 5, p. 222; Cans. Jahresb. d. Pharm., 1862, p. 182.] On the preparation of nitrate and nitrite of ethyl.
The reaction between nitric acid and alcohol in preparation of ethyl nitrite is moderated by ferrous sulphate. "
Those wishing to explore this route should be well acquainted with Ethyl nitrite's MSDS (a link provided) given its poisonous and explosive
properties, and handle only in alcoholic solutions.
http://www.msdshazcom.com/REFERENCE/HAZARDS/ETHYL%20NITRITE....
[Edited on 18-4-2012 by AJKOER]
|
|
|
franklyn
International Hazard
Posts: 3026
Registered: 30-5-2006
Location: Da Big Apple
Member Is Offline
Mood: No Mood
|
|
A lttle paper from 1905 on inorganic nitrites and their decomposition by heat.
How to derive by ion exchange in solution with silver nitrite.
Attachment: Nitrites & decomposition by Heat.pdf (738kB)
This file has been downloaded 791 times
Zinc Nitrite , Zn(NO2)2
CAS 10102-02-0
Literature References: Prepared by treating Sodium Nitrite with Zinc Sulfate in
alcohol: Ephraim, Bolle, Ber. 48, 643 (1915).
Properties: Hydrolyzes so quickly that it cannnot be prepd from water, from
which only basic salts will separate.
Magnesium Nitrite , Mg(NO2)2
is prepared by a solution of Barium Nitrite and Magnesium Sulphate, and has also
been obtained by treating Silver Nitrite with Magnesia or with Magnesium Chloride.
The trihydrate, Mg(NO2)2.3H2O, crystallises in deliquescent leaflets which are
yellowish or snow white. It slowly decomposes in a stoppered bottle, and its
solution decomposes, evolving nitric oxide when evaporated on the water bath.
Careful concentration under diminished pressure , or under diminished pressure
over sulphuric acid , results finally in the dihydrate. This has also been obtained
by digesting magnesium sulphate and sodium nitrite with 94 per cent , alcohol
and evaporating under reduced pressure. It occurs as a hard , efflorescent ,
white mass or as clear crystals , and is much more stable than the trihydrate.
It does not usually dissolve to a clear solution , and may partially decompose
on continued dehydration.
.
[Edited on 15-10-2012 by franklyn]
|
|
|
Natures Natrium
Hazard to Others
Posts: 163
Registered: 22-12-2004
Member Is Offline
Mood: No Mood
|
|
Just wanted to jot down a few quick observations I have made.
10g of KNO3 (directly from stump remover bottle) was held over a propane fueled bunsen burner flame in a 50mL procelain crucible for about 2 hours,
during which time the steel tongs which held the crucible were glowing bright red but the crucible did not appear to be. A test with sulfuric acid
did not indicate the presence of any substantial quantity of nitrite. A comparison with a pre-burn weight showed no detectible (> .1g) weight
loss.
A control test tube with KNO3 had exactly the same non-reaction to sulfuric acid.
Next, 1g of KNO3 in a test tube was slowly warmed at first, and then strongly heated in the flame of a propane blow torch. The salt foamed a bit on
melting, then became water clear. On further heating, when the glass and molted salt were both cherry red, the salt began to evolve gas. This was
not timed, but continued for roughly about 5 to 10 minutes. At this point the glass had sagged slightly due to gravity, and was allowed to cool
slowly by "flame polishing" until all visible glow had ceased. It was then allowed to cool, and after the salt solidified it held the glass together
as the glass in the heating zone cracked into dozens of tiny pieces. A sulfuric acid test was largely positive for nitrites, emitting a good deal of
brown oxides of nitrogen, and most of the salt dissolving.
Interestingly enough, the test tube also tested negative for weight loss, but this may be due to the inaccuracy of the old triple beam to measure
weight differences in the 0.01g to 0.1g range.
Just thought I would share that the thermal decomposition works at least qualitatively, but takes very high temperatures; thus standard borosilicate
is unsuitable for the task.
Edit:Second Attempt:
The same 10g of KNO3 in the same crucible was subjected to a propane bunsen burner flame. This time, the tongs held onto the crucible much further up
the body of the curcible, and the crucible was better situated in the hottest part of the flame (took a bit to relearn my bunsen burner skills). Also
a sheet of aluminum foil in a cone shape was used to help reflect heat inward to the crucible.
After a bit the crucible was glowing orange-red in the dark, and small bubbles could be seen making their way to the surface. It was left in this
state for roughly half an hour, at which point the supposed oxygen evolution had slowed considerably.
At this point the molten salt was poured directly into a clean but old and cracked beaker. After cooling some salt was scrapped up and tested for
nitrite with sulfuric, and the salt tested negative. The beaker (which cracked plenty more) was broken and the salt harvested, several pieces of
which tested negative for nitrite.
Well, I am at a loss. Thermal decomposition of the potassium salt of nitrate clearly works, but also doesn't seem to be very reliable. A reducing
agent seems like it would make things much more dependable. I have some lead shot around here that I may give a try with at some point.
[Edited on 20-2-2013 by Natures Natrium]
\"The man who does not read good books has no advantage over the man who cannot read them.\" - Mark Twain (1835-1910)
|
|
|
AJKOER
Radically Dubious
Posts: 3026
Registered: 7-5-2011
Member Is Offline
Mood: No Mood
|
|
Quote: Originally posted by JohnWW  | If you have a means of generating large amounts of the gaseous lower oxides of N2, particularly NO and N2O, nitrites should be easily obtainable by
bubbling the gas into a concentrated alkali solution.
.......
John W. |
Yes, there is some truth in this statement. Per recent research, apparently in an aqueous environment, the products are either nitrite and nitrate
ions or just nitrite ions depending on pH at room temperature. To quote (source: "Mechanism of the NO2 conversion to NO2- in an alkaline solution" by
Chen X, Okitsu K, Takenaka N, Bandow H. at the Department of Applied Materials Science, Graduate School of Engineering, Osaka Prefecture University):
"The reaction of NO2 and NaOH aqueous solution at room temperature was studied for elucidating the behavior of gaseous NO2 in an alkaline solution.
Experimental runs related to NO2 absorption have been carried out in various pH solutions. The nitrite and nitrate ions formed in these absorption
solutions were quantitatively analyzed. In the case of pH 5-12, both of the nitrite and nitrate ions were formed simultaneously. On the other hand,
only the nitrite ion was formed when the pH of the absorption solution was higher than 13. In this paper, a new reaction mechanism was proposed to
explain the selective formation of nitrite ion in the 10 M alkaline solution. In order to confirm the new reaction mechanism, H2(18)O was used as part
of the absorption solution for detecting oxygen gas production. The amounts of reaction products: (18)O(18)O, (18)O(16)O and (16)O(16)O, were
quantitatively determined. It was confirmed that the new reaction proceeds mainly in the 10 M alkaline solution."
Link: http://www.ncbi.nlm.nih.gov/pubmed/15636532
Now, if solid moist NaOH is treated with NO2, in this high pH condition, I suspect that primarily NaNO2 may form. However, this study using Soda Lime
("Elimination of nitrogen dioxide and nitric oxide by soda lime and its components", original in Chinese by Zhang D, Hu X, Liu J at the Department of
Anesthesiology, Fu Wai Hospital, Chinese Academy of Medical Sciences, Peking Union Medical College) for scrubbing gases notes that NaOH or KOH by
themselves are not effective in removing NO and NO2 unless Ca(OH)2 is present. Link: http://www.ncbi.nlm.nih.gov/pubmed/9772493
As such, I would also consider adding Ca(OH)2 to increase reactivity between the moist NaOH and gaseous NO2 at a high pH. I would expect the primary
product to be NaNO2 (and some Ca(NO2)2). My take on the reaction:
4 NaOH + 4 NO2 --High pH--> 4 NaNO2 + 2 H2O + O2
|
|
|
jock88
National Hazard
Posts: 505
Registered: 13-12-2012
Member Is Offline
Mood: No Mood
|
|
Some good info. in this thread about using Tin instead of lead
for making nitrites.
http://www.sciencemadness.org/talk/viewthread.php?tid=23485
I guess a mod should merge threads??
|
|
|
AJKOER
Radically Dubious
Posts: 3026
Registered: 7-5-2011
Member Is Offline
Mood: No Mood
|
|
Here is a commercial based method, using Cadmium (and Cadmium treated with Copper) that is complexed either by NH4Cl or the Sodium salt of EDTA. In a
neutral to basic solution the reaction is given by:
[NO3]- + H2O + 2 e- ---> [NO2]- + 2 [OH]-
and the oxidation of Cadmium by:
Cd + 1/2 O2 + H2O --> Cd(OH)2
or, if using EDTA:
[NO3]- + Cd + [EDTA]4- + H2O ---> [NO2]- + Cd[EDTA]2- + 2 [OH]-
The method is claimed to have near quantitative reduction of nitrate to nitrite.
Source: Journal of the Marine Biological Association of the United Kingdom / Volume 47 / Issue 01 / February 1967, pp 23-31. Link to full text: http://www.google.com/url?sa=t&rct=j&q=determination...
For the current application of processing large amounts of NaNO3, I would start by dissolving NaNO3 in hot water and add powdered Cd, and further
heat. One could improve the reactivity of the Cadmium by employing the thermal decomposition product of Cd oxalate in nitrogen. To quote: "It is
suggested that the decomposition of cadmium oxalate to cadmium and cadmium oxide is ... and stated that the formation of metal is the result of
reduction of oxide". A reference link to the entire article (for a fee) can be found at http://www.sciencedirect.com/science/article/pii/00406031828... . I would prefer this route over complexing with NH4Cl as any excess Ammonium
chloride could act on the newly created NaNO2 as follows reducing yield:
NH4Cl + NaNO2 --> NaCl + NH4NO2
as the Ammonium nitrite formed is unstable decomposing into water and N2 (and in concentrated/acidic solutions in an energetic manner).
[Edited on 6-4-2013 by AJKOER]
|
|
|
AJKOER
Radically Dubious
Posts: 3026
Registered: 7-5-2011
Member Is Offline
Mood: No Mood
|
|
For those wishing to work with more friendly metals and readily available to turn nitrates into nitrites, consider Aluminum (Al foil no less). Per one
source (see equation [1] under Section 6.4, link:
https://docs.google.com/viewer?a=v&q=cache:Jz8OCxPNXSoJ:...), to quote:
"1. 3NO3- + 2Al + 3H2O → 3NO-2 + 2Al(OH)3
2. NO-2 + 2Al + 3H2O → NH3 + 2Al(OH)3 + OH-
3. 2NO-2 + 2Al + 4H2O → N2 + 2Al(OH)3 +2OH-
Nitrate reduction was found to be pH dependant. At pH values less than eight
no nitrate reduction took place. Above pH 10.5 nitrate was reduced upon addition of
the aluminium powder. Aluminium powder has been suggested for the denitrification
of sodium-based nuclear wastes, employing the nitrate to ammonia and ceramic
(NAC) process (Mattus et al., 1993, 1994)."
Here are encouraging comments from another source (see http://en.wikibooks.org/wiki/Inorganic_Chemistry/Qualitative... ):
"The Nitrate ion can easily be reduced to ammonia with either Devardas Alloy or Aluminium Foil. The aluminium is a very powerful reducing agent, and
this combined with heating causes the nitrate ions to form ammonia gas. This can be tested for by holding a piece of damp red litmus paper over the
end of the test tube. The ammonia will form alkaline ammonium ions in the water and turn the paper blue.
4NO3-(aq) + 6H20(l) -> 4NH3(g) + 9O2(g)
Aluminium powder is not shown as it merely catalyses the reaction."
As a third, and some would claim as less credible authority, I present my analysis on the mechanism on exactly how Aluminum is involved in reaction
[1]. First, as the Al only initiates the reaction in only more alkali conditions, I suspect the following occurring:
2 Al + 6 H2O --High pH Weakening Al2O3 Coating--> 2 Al(OH)3 + 3 H2 (g)
Or, more directly per the action of the strong alkali, for example;
2Al(s) + 2NaOH + 6H2O → 2Na[Al(OH)4] + 3H2(g)
That is, I would claim, it is nascent Hydrogen (label it H*) responsible and as some support, I do recall reading in an old text that freshly prepared
Hydrogen can reduce nitrates (I look for the reference). So, a possible reaction, for example:
3 NaNO3 + 6 H* --> 3 NaNO2 + 3 H2O
So, in my opinion, it is the nascent H2, and not directly the Aluminum, that is the so called powerful reducing agent here. Following this point,
nascent Hydrogen generation say from an external metal/acid reaction (like Zn/HCl) may also prove to be successful. Here is a source on nascent H2
(see http://en.wikipedia.org/wiki/Nascent_hydrogen ), to quote:
"According to one claim, nascent hydrogen is generated in situ usually by the reaction of zinc with an acid, aluminium (Devarda's alloy) with sodium
hydroxide, or by electrolysis at the cathode.[citation needed] Being monoatomic, H atoms are much more reactive and thus a much more effective
reducing agent than ordinary diatomic H2, but again the key question is whether H atoms exist in any chemically meaningful way under the conditions
claimed. The concept is more popular in engineering and in older literature on catalysis."
A second, and very important point, is that per the three equations given above, and the evidence of ammonia formation per the 2nd author, one cannot
employ an excess of Aluminium (else further reduction to NH3 and N2 could occur), and also, I would recommend stirring. The good news is that the
reaction proceeds on gentle heating.
[EDIT] I found an old reference, "A manual of chemical analysis, qualitative and quantitative", by George Samuel Newth, page 146, link: http://books.google.com/books?pg=PA146&lpg=PA146&dq=... ) to quote:
"Reduction by Nascent Hydrogen. (1) With Formation of Nitrite.—When a nitrate in solution is exposed to the gentle action of nascent
hydrogen—derived by the action of sodium amalgam, zinc amalgam, or copper-zinc couple—the nitrate is reduced to nitrite—
KN03 + H2 = H20 + KN02
This test is very delicate, and may be carried out as follows: A small piece of zinc foil (or granulated zinc) is placed in the solution of the
nitrate in a test-tube, and one drop of copper sulphate added (this causes the deposition of a minute quantity of copper upon the zinc, thus creating
the "copper-zinc couple"). The mixture is gently boiled for a minute or two. 0ne drop of the liquid (after cooling) is placed upon a piece of
potassium-iodide-and-starch paper, and then touched with a glass rod moistened with dilute sulphuric acid. The paper will be instantly stained blue
by the liberation of iodine and formation of iodide of starch (test for a nitrite).
[foot note]
* The mechanism of this reaction is sometimes explained by supposing that the nitric acid, set free from the nitrate, acts upon the copper according
to the familiar equation 3Cu + 8HN03 = 3Cu(N03)2 + 4H20 + 2N0. But, as a matter of fact, copper sulphate, not nitrate, is found in solution, the whole
of the nitrogen being converted into nitric oxide. Moreover, charcoal may be substituted for the copper. If dilute nitric acid be boiled with charcoal
no brown fumes are formed, but on the addition of a little sulphuric acid they at once appear, owing to the action of the sulphur dioxide which is
evolved from carbon and sulphuric acid."
My new take on the reputed creation of nascent H2 by the several mentioned paths is that it appears to me that there may be, in fact, an
electrochemical connection (note, my quote source above on nascent Hydrogen citing one possible formation also "by electrolysis at the cathode").
Also, chemically pure Aluminum (or Zinc) may be inferior to say Al foil, an amalgam.
A note on Devarda's alloy (aluminium (44% – 46%), copper (49% – 51%) and zinc (4% – 6%)), to quote (see
http://en.wikipedia.org/wiki/Devarda's_alloy ):
"Devarda's alloy is used as reducing agent in analytical chemistry for the determination of nitrates after their reduction to ammonia under alkaline
conditions. It owes its name to the Italian chemist Arturo Devarda (1859–1944), who synthezised it at the end of the 19th century to develop a new
method to analyze nitrate in Chile saltpeter.[2][3][4]"
Also, "The reduction of nitrate by the Devarda's alloy is given by the following equation:
3 NO3− + 8 Al + 5 OH− + 18 H2O → 3 NH3 + 8 [Al(OH)4]− "
This reaction with excess Aluminum is in complete agreement with a prior comment by Woelen with a qualification that I am suggesting by limiting the
Al, per equation [1] above, and per an old qualitative test for nitrite, one may be able to successfully reduce nitrates to nitrite. To quote:
| Quote: Originally posted by woelen | Trying to use iron in aqueous solution will not produce any nitrite. Nitrate ion is remarkably inert in aqueous solution, except in strongly
concentrated acidic environments or at very high pH in combination with magnesium and aluminium. In the acidic environments it forms NOx with suitable
reductors, in the strongly alkaline environment it forms ammonia.
..... |
[Edited on 8-4-2013 by AJKOER]
|
|
|
Oscilllator
National Hazard
Posts: 659
Registered: 8-10-2012
Location: The aqueous layer
Member Is Offline
Mood: No Mood
|
|
A couple of my own observations regarding potassium nitrate -> nitrite through thermal decomposition:
Over several batches, approx 250g KNO3 was boiled in a stainless steel saucepan for about 15 minutes. The source of the heat was a coke-fired forge
that had previously proved itself capable of melting steel During the boiling, copious quantities of smoke was evolved.
When the molten salt was cooled down, it formed a distinctly green solid (greener than it looks in the photo). perhaps this is some kind of iron
compound from the stainless steel?
This was then dissolved in approx 500ml of water, heated by placing it on the forge. This brought the 500ml of water to the boil in about 10-15
seconds  . The resultant brown muddy liquid was filtered to obtain a clear
yellow solution that formed brown fumes upon addition of HCl. When the liquid was cooled down, large quantities of needle-shaped crystals formed that
proved to be potassium nitrate  .
This solution was then boiled down and a second crop of crystals were formed. They had a slightly different shape, and I am not certain that they are
KNO3 or KNO2, but I'm afraid the odds favour the KNO3. Irrespective of that, large quantities of KNO2 were formed in this process, as least enough
that the evolution of NO2 was so great that the test tube almost bubbled over.
It is interesting to note that wikipedia and a number of other sources say that KNO2 explodes when heated above around 700 degrees. I am 90% sure this
is not the case based on the fact that a rounded tablespoon of KNO3 was fully decomposed all the way to some black gunk that bubbled when water was
dropped on it. Presumably this black substance is K2O. At any rate, it was sitting on top of a red-hot piece of steel and didn't do anything that
could be described as an explosion  .
I do have a hypothesis regarding the entire method of producing nitrites by thermal decomposition: Is it possible that as the nitrate decomposes the
formed nitrite also decomposes, so at no point in time is there a solution (if thats the right word) of 100% nitrate?
|
|
|
plante1999
International Hazard
Posts: 1936
Registered: 27-12-2010
Member Is Offline
Mood: Mad as a hatter
|
|
Nice try, but it was known thermal process works so so and need multiple fractional crystallization. Try to use lead or copper powder next time.
I never asked for this.
|
|
|
S.C. Wack
bibliomaster
Posts: 2419
Registered: 7-5-2004
Location: Cornworld, Central USA
Member Is Offline
Mood: Enhanced
|
|
| Quote: Originally posted by Oscilllator |
It is interesting to note that wikipedia and a number of other sources say that KNO2 explodes when heated above around 700 degrees.
|
Boiling point 537 °C (explodes) may be total bullshit. Perhaps 700C is what it takes for this to "work" (15 minutes is probably still not
long enough), as has been noted in this thread.
http://www.sciencemadness.org/talk/viewthread.php?tid=52&...
PS My summary of all this is: probably best done with some form of protection from O at a temperature not much below where N in some form is lost. No
explosions IME.
[Edited on 4-8-2013 by S.C. Wack]
|
|
|
watson.fawkes
International Hazard
Posts: 2793
Registered: 16-8-2008
Member Is Offline
Mood: No Mood
|
|
| Quote: Originally posted by Oscilllator | When the molten salt was cooled down, it formed a distinctly green solid (greener than it looks in the photo). perhaps this is some kind of iron
compound from the stainless steel?
|
Chromium oxide, I'd guess. Stainless steel is stainless because it forms passivating surface coatings of
chromium oxides. I'd guess you've liberated some of the Cr from the outer layer of the pot.
|
|
|
blogfast25
International Hazard
Posts: 10562
Registered: 3-2-2008
Location: Neverland
Member Is Offline
Mood: No Mood
|
|
| Quote: Originally posted by watson.fawkes | | Chromium oxide, I'd guess. Stainless steel is stainless because it forms passivating surface coatings of chromium oxides. I'd guess you've liberated
some of the Cr from the outer layer of the pot. |
Yes, that happens each time I use a SS 'crucible' with some alkaline fusion. The iron itself doesn't seem to be affected much but Cr is amphoteric.
|
|
|
AJKOER
Radically Dubious
Posts: 3026
Registered: 7-5-2011
Member Is Offline
Mood: No Mood
|
|
| Quote: Originally posted by AJKOER | For those wishing to work with more friendly metals and readily available to turn nitrates into nitrites, consider Aluminum (Al foil no less). Per one
source (see equation [1] under Section 6.4, link:
https://docs.google.com/viewer?a=v&q=cache:Jz8OCxPNXSoJ:...), to quote:
"1. 3NO3- + 2Al + 3H2O → 3NO-2 + 2Al(OH)3
2. NO-2 + 2Al + 3H2O → NH3 + 2Al(OH)3 + OH-
3. 2NO-2 + 2Al + 4H2O → N2 + 2Al(OH)3 +2OH-
Nitrate reduction was found to be pH dependant. At pH values less than eight
no nitrate reduction took place. Above pH 10.5 nitrate was reduced upon addition of
the aluminium powder. Aluminium powder has been suggested for the denitrification
of sodium-based nuclear wastes, employing the nitrate to ammonia and ceramic
(NAC) process (Mattus et al., 1993, 1994)."
.........
|
A week ago, I tried to convert some of my KNO3 to KNO2 using Al foil per Equation [1]. I am now reporting some of the results, for those interested,
to the best of my recollection.
I restricted the aluminum to the amount required for just Equation [1] per above:
3NO3- + 2Al + 3H2O → 3NO-2 + 2Al(OH)3
and also added Na2CO3 to make the solution alkaline as was required per the author's discussion. After the 1st day, visible signs of Al(OH)3. I
extract some of the solution to experiment further with the nitrite (another story).
What is interesting is what happened in the remaining solution in the following day. A significant gas formation became evident. When I opened the
vessel, I was greeted by a rush of ammonia fumes. This was in accord with Equation [2] above, namely:
NO-2 + 2Al + 3H2O → NH3 + 2Al(OH)3 + OH-
My conclusion, was that removing of the some of aqueous KNO3/KNO2 may have created an excess of Aluminum which relatively rapidly (as compared to Eq
[1]) formed the NH3 per above.
If I repeat the experiment again, I will certainly use an excess of KNO3. Clearly, the significant formation of ammonia implies that the nitrate can
be converted to nitrite and then ammonia by this method. I am not certain, however, as to precise yield measures on the KNO2.
The purpose of my experiment was to test the possibility of nitrite only formation using Aluminum foil. I am more confident now that it may be
possible.
[Edited on 25-12-2013 by AJKOER]
|
|
|
thebean
Hazard to Others
Posts: 116
Registered: 26-9-2013
Location: Minnesota
Member Is Offline
Mood: Deprotonated
|
|
I had a semi successful non stoichiometric micro scale reaction. I put some potassium nitrate in a test tube with iron wool and heated until bubbles
began forming. Small isolated incidents of ignition occurred in the iron but I would remove heating to stop the combustion. After the reaction stopped
I was left with yellowish crystals which I suspect were nitrite. To test for presence of nitrite, I added some hydrochloric acid which produced brown
fumes of what can only be NOx. I wafted a little and it had the classic NOx scent with a hint of chlorine. If one had proper amounts of nitrate and
iron (preferably powder) and placed it in a small FBF with gentle heating up to the melting point of the nitrate and maintain that temperature until
all of the iron appeared to have reacted you could produce decent amounts of the corresponding alkali nitrite, of course I would guess that an order
of nitrite would be cheaper. It appears that you can replace the lead in the typical reaction with a variety of metals, the easiest to acquire being
iron but lead having the best results.
"You need a little bit of insanity to do great things."
-Henry Rollins
|
|
|
AJKOER
Radically Dubious
Posts: 3026
Registered: 7-5-2011
Member Is Offline
Mood: No Mood
|
|
Here is a source confirming the use of a variety of metals. To quote (see https://www.google.com/url?sa=t&rct=j&q=&esrc=s&... ):
"Potassium nitrate decomposes to the nitrite when heated to its melting point. Reducing
agents can also be added to facilitate the reaction to completion these are Pb, Fe2S, Fe2O, Cu, MnO2,
Fe, chromic, and sulfur. Best results were from MnO2 at 650 ° C.
The nitrite can also be formed by reduction of a nitrate with a metal in their aqueous
solutions. Examples of the reducing agents are Zn, K, Na, Pb, or Zn-Amalgam. The problem with
these reactions is the nitrite when heated in the aqueous solution will absorb oxygen from the air
and water, which reforms the nitrate. The electrolysis of sodium nitrate with amalgated copper
cathode. Silver nitrate solution reacted with silver and heated yields silver nitrite.
Ammonia can be oxidized to ammonium nitrite by the addition of hydrogen peroxide or
potassium permanganate.
In the reaction of a nitrate with granulated lead, the reducer should be in 15% excess. The
reaction commences at the boiling point and after 2 hours at 420 °C, complete reduction is obtained
(example with sodium nitrate).
KNO3 + H3PO3 → KNO2 + H3PO4
NaNO3 + CaO + SO2 → CaSO4 + NaNO3
2KNO3 → 2KNO2 + O2 with heat
NaNO3 + R → NaNO2 + RO where R is the metal reducing agent mentioned above
Pb + NaNO3 → PbO + NaNO2 "
--------------------------------------
Here is another old source, "Journal of the Society of Chemical Industry", Volume 27, based on commercial processes with some very insightful
commentary on processes (see pages 484 to 485, link: http://books.google.com/books?pg=PA484&lpg=PA484&dq=... ). For example:
"Thermal decomposition.—The alkali nitrates when heated above their fusion point evolve oxygen and furnish nitrites, but this reaction is, of no
practical importance owing to the simultaneous occurrence of a further decomposition to oxide. "
There is also an interesting aqueous method mentioned. To quote:
"The heating of a concentrated solution of lead nitrate with finely divided lead leads to the formation of insoluble basic lead nitrite, which
furnishes sodium nitrite by double decomposition with sodium carbonate. This reaction is, however, only of theoretical interest. "
Per the prior source, one can employ Zinc in place of Pb in this aqueous reaction. Apparently, Zn dust is effective in reducing nitrate to nitrite
(see, for example, http://www.microbelibrary.org/library/laboratory-test/3660-n... ). To quote:
"FIG. 5. Zinc dust will reduce nitrate to nitrite, but will not further reduce nitrite to nitrogen gas or other nitrogenous by-products when used
sparingly."
The author also notes, to quote:
"3. Some authors recommend adding zinc to colorless NO2- reactions that do not contain gas to make sure that the NO2 has not been oxidized to NO3
rather than having been reduced to a nitrogen product other than N2 gas (21), but that reaction is rare."
which suggests to me the need to remove O2 from the concentrated aqueous nitrate solution by pre-boiling and avoid further exposure to air (perhaps by
adding CO2,..). In place of Zn dust, I might use (when I perform this reaction) a Zn colloidal suspension in an O2 free solution (say from the action
of Aluminum foil on an aqueous Zinc ammonium salt).
[Edited on 5-1-2014 by AJKOER]
|
|
|
Alyosha Karamazov
Harmless
Posts: 4
Registered: 19-10-2013
Member Is Offline
Mood: counterbalancing
|
|
The J. Soc. Chem. Ind. (1908) review on nitrite production mentioned by AJKOER is a very interesting read. Unfortunately I can't view it using google
books in my country, but archive.org is more lenient when it comes to copyright restrictions, I downloaded the journal issue from there and snipped
the article so I could provide it here.
Also included is the Compt. Rend. article from 1889 (and not 1900 stated in the review) describing briefly the reduction of sodium nitrate by heating
with barium sulfide, the mixture diluted with barium sulfate to make it less violent and more controllable. Separation of the nitrite is easy and the
formed BaSO4 could be used to produce the sulfide again, as suggested by the author.
Also provided here is a BASF patent (FR363643) describing the oxidation of nitrogen in air by electrical discharge. Interesting is that the author
notes that while the nitric oxide produced readily oxidises in air to nitrous anhydride (N2O3), the latter oxidises much less
readily to nitrogen dioxide at elevated temperatures. The rate of oxidation of N2O3 to NO2 is 20x higher at 0°C
compared to 100°C, and at 300°C it is virtually non-existant. Hence the author suggests working at that temperature to transfer and absorb the
nitrous anhydride in basic solutions.
Attachment: J.Soc.Chem.Ind.27_1908-nitrite_review.pdf (491kB)
This file has been downloaded 638 times
Attachment: FR321498A-reduction_using_finely_divided_metals.pdf (86kB)
This file has been downloaded 671 times
Attachment: CR_1889_108-reduction_using_barium_sulfide.pdf (211kB)
This file has been downloaded 646 times
Attachment: FR363643A-electrical_discharge_absorption_of_N2O3.pdf (90kB)
This file has been downloaded 676 times
|
|
|
S.C. Wack
bibliomaster
Posts: 2419
Registered: 7-5-2004
Location: Cornworld, Central USA
Member Is Offline
Mood: Enhanced
|
|
| Quote: Originally posted by komodo13 | | J. Soc. Chem Ind., 27, 483-5 (may30)-The author reviews various methods for the production of nitrite from nitrate by reduction with metals
|
Yesterday I also happened to independently find this article while searching for something totally unrelated. Searching the author's name just now
shows that just a couple days before that, a different article by the same author was requested in the references thread. And, this person here who
signed up a month after me had come by to drop the name, made one more post, then disappeared. I'm not sure if posting the full text, and the author's
name for the third time onsite, is the last thing necessary to start the Apocalypse or what, but now I'm afraid not to. I better include the abstract
of his obit in a 1940 issue of Nature: SIR GILBERT MORGAN died after a very short illness on February 2 in his seventieth year. For fifty years he
had been engaged in chemical research and probably no other chemist had such wide knowledge of scientific and applied chemistry. A fitting mark of the
Jubilee was the presentation to him in July last of the medal of the Society of Chemical Industry, its highest honour; on this occasion he gave an
account of his career and an outline of his researches.
THE MANUFACTURE OF SODIUM NITRITE.
GILBERT T. MORGAN
J. Soc. Chem. Ind. 27, 483 (1908)
Sodium nitrite is practically the only salt of nitrous acid which is prepared on a manufacturing scale and it finds extensive use in the production of
several classes of artificial colouring matters and also in the preparation of various pharmaceutical products and other fine chemicals.
I. Production from sodium nitrate.
Hitherto nitrite has generally been prepared from sodium nitrate, and the following are some of the principal methods by which this chemical change
can be effected.
1. Thermal decomposition.—The alkali nitrates when heated above their fusion point evolve oxygen and furnish nitrites, but this reaction is of no
practical importance owing to the simultaneous occurrence of a further decomposition to oxide. The difference in the behaviour of the two alkali
nitrates is of some theoretical interest, the potassium salt giving a final residue of the dioxide, whilst the sodium salt yields the monoxide.
2. Reduction by metals.—Finely divided copper has been suggested by Persoz and by Muller and Pauly, but owing to the infusibility of this metal at
the temperature of molten nitre it is difficult to ensure a uniform reduction throughout the heated materials, and, moreover, this process involves
the regeneration of the comparatively expensive copper from the resulting copric oxide. Zinc dust has been tried but I am not aware that a successful
method has been based on the use of this metal. Sturm (Fr. Pat. 321,498 of 1902; this J., 1903, 212) obtains nitrite by heating sodium nitrate with
finely divided metals in a muffle furnace. More recently iron has been recommended as a reducing agent for sodium nitrate.
The heating of a concentrated solution of lead nitrate with finely divided lead leads to the formation of insoluble basic lead nitrite, which
furnishes sodium nitrite by double decomposition with sodium carbonate. This reaction is, however, only of theoretical interest.
Lead is undoubtedly the most convenient metal for the reduction of sodium nitrate; it is comparatively cheap, and its melting point (330—335° C.)
lies so close to that of the nitrate (314° C.), that the two reagents can be brought together in the liquid condition at a temperature considerably
lower than that at which the thermal decomposition of nitrite occurs. (Hampe, Annalen, 1863, 125, 336.) The reduction is effected on a manufacturing
scale in shallow cast-iron pans, 4 ft. in diameter and about 18 inches deep, fitted with a stirring gear which agitates thoroughly the whole mass of
molten material. The pans are supported on perforated firebrick arches, arranged so that the products of combustion of the coal fires circulate
uniformly round the pan before passing to the fines. Each pan is charged with 200 lbs. of sodium nitrate and 3 cwts. of lead, which are heated and
stirred until all the metal is oxidised. At this stage 50 lbs. of nitre are added and thoroughly stirred in until the mixture is of uniform
consistence when 3 1/2 cwts. of lead are gradually added, with constant agitation, the mechanical stirring being supplemented by the use of
long-handled rakes employed to “pull out" the mixture of metal and salt from the central and hottest part of the pan. The success of the reduction
depends very largely on the skill and experience of the workman, who generally controls two pans, and whose duties are to regulate the firing of the
pan and the rate of addition of the metal.
The reaction occurring in the nitrite pan is not so simple as that represented bv the equation, Pb + NaNO3 = PbO + NaNO2, for the higher oxides of
lead are also produced both by aerial oxidation and by further reaction with the nitrate. In this connection it should be noted that litharge itself
has been suggested as a reducing agent for nitrate, 3PbO + NaNO3 = NaNO2 + Pb3O4, although the large proportion of this oxide required would militate
against its adoption.
During the reaction a portion of the lead becomes converted into a singularly inert substance of high specific gravity; this product is regarded as a
sub-oxide, but may be a passive form of the metal contaminated with higher oxides and sodium plumbite.
The fusions are allowed to run for one-half to three quarters of an hour after all the lead has been added, when the product, which now has a
yellowish brown tint, is tested for nitrite. With careful working the soluble constituent of the melt should contain 90 per cent. or even more of
sodium nitrite. When rich in nitrite the cooled melt has on its surface a characteristic crystalline incrustation which is never noticed on specimens
containing a relatively small percentage of the required salt.
Satisfactory results are, however, only obtained when both the lead and nitrate are of good quality. The former should be good commercial lead which
has been remelted, skimmed and cast into small bars. The latter should be crystallised Chili saltpetre of the best quality. The presence of sodium
iodate is especially harmful, as this salt appears to act catalytically in promoting the destruction of the molten nitrate and nitrite. The appearance
of the characteristic violet vapour of iodine arising from the melt shows that inferior nitrate is being employed and in these circumstances the salt
should be recrystallised, when the harmful impurities are eliminated.
The reduction being complete, the molten contents of two pans are ladled into 120 gallons of warm water contained in a covered cylindrical washing box
fitted with powerful stirrers rotating on a horizontal axis. The curved lid of this box contains two small circular apertures for the introduction of
the melt, and within, the soluble nitrite is separated from the litharge by agitating the mixture for one hour and then allowing the precipitate to
subside. The clear liquor is then run off into a neutralising tank and here the solution, which is distinctly alkaline, is neutralised with dilute
sulphuric acid or, if possible, with the solution obtained by absorbing nitrous fumes in water. These fumes are obtained in such nitric acid oxidation
processes as the manufacture of arsenic acid.
This neutralisation causes the decomposition of sodium plumbite and the precipitation of a "small amount" of lead hydroxide. The neutralised liquors
are now evaporated in wrought iron pans heated either directly or with internal steam pipes. At 45° Bé. the concentrated solution shows a thin film
of nitrite on its surface and is then run off into rectangular cast iron crystallisers and left for at least 12 hours, when the first crop of sodium
nitrite crystals is collected, the mother liquors being again concentrated and allowed to crystallise. These crystals, when dried in a centrifugal
hydro-extractor and then in an air-oven, should contain 96 per cent. of NaNO2.
The washing of the oxides of lead is repeated, the more dilute washing liquors being flushed off together with the litharge into settling tanks. These
liquors are used to lixiviate subsequent melts, whilst the litharge is either dried for sale or mixed with the “sub-oxide” and smelted to lead in
a small blast furnace.
As received from the settling tanks the pasty litharge contains varying proportions of the higher oxides of lead. It may, however, be rendered more
uniform by conversion into flake litharge or red lead in suitable reverberatory furnaces.
The working up of these large quantities of lead compounds is one of the chief disadvantages of this process, another is the baleful effect of the
lead on the workmen, a certain incidence of plumbism being almost unavoidable.
3. Reduction by non-metals.—The well-known detonation of charcoal and nitre leads to the formation of a carbonate and only small quantities of
nitrite. When brought under control by the addition of caustic soda and lime the reaction between sodium nitrate and graphite has been patented as a
process for preparing sodium nitrite (Grossmann, Eng. Pat. 1452 of 1904; and also Knop, Eng. Pat. 4747 of 1897).
A similar process involving the use of sulphur, which has been successfully worked out on a manufacturing scale by Messrs. Read Holliday and Sons, of
Huddersfield, is based on the following reaction:—
3NaNO3 + S + 2NaOH = Na2SO4 + 3NaNO2 + H2O.
The fusion is carried out in open pans fitted with stirring gear, but of larger capacity than those employed in the lead process. The nitrate
containing a portion of the caustic soda (about 1/10th) is melted and treated alternately with sulphur and more molten caustic soda until the nitrate
is practically all reduced. The fused product is added while still hot to sufficient warm water to dissolve the whole of the nitrite and only a
portion of the sodium sulphate, a large proportion of which is left behind in a granular condition. The liquor is drained through a vacuum filter and
evaporated to a smaller bulk when a further portion of sulphate separates and the solution on cooling deposits sodium nitrite while the final mother
liquors furnish more Glauber’s salt.
This process yields without troublesome by-products a nitrite of good quality, which is obviously quite free from lead, and apart from the separation
by fractional crystallisation of the nitrite from the dissolved sulphate, the operations involved present no serious practical difficulties.
Carbon monoxide either pure or in the form of producer gas has no action on molten sodium nitrate but in the presence of fused caustic soda there in
an intermediate formation of formate which then reduces the nitrate in the following manner:—
NaNO3 +HCO2Na + NaOH = NaNO2 + Na2CO3 + H2O.
(Goldschmidt, Eng. Pat. 17066 of 1895.)
4. Reduction by metallic sulphides and sulphites.—Etard formerly recommended sodium sulphite as a reducing agent for sodium nitrate (Bull. Soc.
Chim., 1877, 27, 434), and a modification of his process has recently been patented (D.R.P. 138,029). The percentage yield of nitrite is excellent but
the high proportion of sodium sulphate in the melt - about two-thirds of the total — renders the separation of sodium nitrite somewhat troublesome.
Closely allied to this method is the process devised by the firm of Gebruder Flick, which consists in passing sulphur dioxide over sodium nitrate and
calcium hydroxide heated in retorts. The nitrite is then readily separated from the sparingly soluble calcium sulphate.
The commonly occurring sulphides react with fused sodium nitrate, furnishing nitrite. Messrs. McGougan have patented the use of galena which gives a
melt containing litharge, sodium nitrite and sulphate with a small amount of sodium plumbite (Eng. Pat. 7715 of 1897). I have noticed that stibnite,
the fusible sulphide of antimony, very readily reduces the nitrate giving a high percentage of nitrite but the cost of this reducing reagent is
prohibitive.
Sodium sulphide and nitrate interact energetically forming a nitrite melt which contains only a relatively small proportion of sodium sulphate. Le Roy
has advocated the use of barium sulphide, a mixture of this substance and sodium nitrate being heated in an iron dish when a vigorous reaction sets in
and sodium nitrite and barium sulphate result. The intensity of this reduction is moderated by the admixture of barium sulphate (Compt. rend., 1900,
108, 1251).
II. Production of nitrite from nitrous fumes.
It was shown conclusively by Divers (Trans., 1899, 75, 85) that pure sodium nitrite could be readily prepared by absorbing nitrous fumes in aqueous
sodium carbonate or hydroxide, provided that these gases contain a slight excess of nitric oxide. Excess of nitrogen peroxide would result in the
formation of nitrate. Nitric oxide itself was shown by Debray to unite with barium peroxide forming barium nitrite, a similar reaction with the alkali
peroxides would lead to sodium and potassium nitrites.
Raschig‘s observation that nitric oxide combines very rapidly with oxygen to form nitrous anhydride whilst the further change of the latter oxide to
nitrogen peroxide occurs comparatively slowly suggests a method of utilising atmospheric nitrogen in the production of nitrite. The absorption of the
nitrous fumes within a few seconds of their formation in the electric arc is an operation involving considerable practical difficulty, which however
has to some extent been overcome by the method recently patented by the Badische Anilin- und Soda-fabrik (Fr. Pat. 363,643 Of 1906). According to this
patent the nitrous fumes are maintained at a temperature of 300° C. until they are absorbed by an alkaline solution of sodium nitrite from a former
operation. A strong solution of this salt is employed in order to reduce as far as possible the vapour pressure of the liquid, and thus minimise the
dilution of the hot reacting gases with steam.
According to Eyde (Eng. Pat. 28,613 of 1904) the gases of the electric furnace containing much air when quickly brought into contact with the
hydroxides of the alkalis or alkaline earths, yield nitrites
2NO + NaOH + O = 2NaNO2 + H2O.
Electrolytic reduction of nitrates.-Various attempts have been made to utilise the electric current in the reduction of nitrates to nitrites. Among
the most recent are the experiments made by E. Muller and F. Spitzer (Zeit. Elektrochem., 1905, 11, 509) with cathodes of different metals, the most
favourable results being obtained with spongy silver.
Miscellaneous agents.-The interaction between barium hydroxide, manganese dioxide and sodium nitrate has been patented by Huggenberg. Zinc and ammonia
have been employed by Stahlschmidt. The oxidation of ammonia in the presence of metallic oxides at 650-750° C. leads to the production of “nitrous
anhydride” which is absorbed by alkalis (U.S. Pat. 763,491 of 1904). This oxidation of ammonia to nitrite has also been effected electrolytically in
aqueous solution in the presence of sodium and cupric hydroxides (Traube and Biltz, Ber., 1904, 37, 3120).
III. Production of nitrite from calcium nitrate.
As it has been predicted that calcium nitrate will gradually displace the sodium salt as the commercial source of nitre, I have made some experiments
on the production of sodium nitrite from calcium nitrate, or from mixtures of this salt with calcium nitrite.
Calcium nitrate melts in its water of crystallisation becomes solid again at higher temperatures and finally fuses. When maintained in a pasty state
for some time the anhydrous salt loses oxygen and oxides of nitrogen; some nitrite is produced but the yield is very small. Reducing agents increase
the production of nitrite very considerably and when mixed with sodium sulphite and sulphide the calcium nitrate on heating furnishes a yield of more
than 60 per cent. of the calculated amount of sodium nitrite. The object of taking the two reducing agents in these proportions is to ensure the
conversion of both sulphur compounds into sparingly soluble calcium sulphate.
i. 2Ca(NO3)2 + Na2S = CaSO4 + Ca(NO2)2 + 2NaNO2
ii. Ca(NO3)2 + 2Na2SO3 = CaSO4 + Na2SO4 + 2NaNO2
The combined changes may be represented as follows:-
3Ca(NO3)2 + Na2S + 2Na2SO3 = 6NaNO2 = 3CaSO4
By taking the sulphide and sulphite in these proportions the product after lixiviation consists chiefly of very soluble nitrite and sparingly soluble
gypsum which are readily separated. The sulphide and sulphite are melted together until their water of crystallisation is driven off and the residue
intimately mixed with the calcium nitrate. This mixture is heated until the water contained in the last salt is eliminated. A portion of the dried
mixture is then heated strongly until a reaction sets in accompanied by incandescence and the remainder is added sufficiently rapidly to ensure the
continuance of this interaction. The greyish white product is lixiviated with warm water, the sulphate removed and the nitrite obtained from the
solution.
With a mixture of nitrate and nitrite the sulphide may be omitted as in this case sulphite alone suffices to convert all the calcium into sulphate.
Ca(NO2)2 + Ca(NO3)2 + 2Na2SO3 = 4NaNO2 + 2CaSO4.
Instead of the sulphide and sulphite, a mixture of sulphur, caustic soda and calcium nitrate may be employed.
3Ca(NO3)2 + 2S + 6NaOH = 2CaSO4 + Ca(OH)2 + 6NaNO2.
The introduction of carbon dioxide or dilute sulphuric acid into the aqueous solution of the melt ensures the precipitation of the calcium hydroxide
in the form of calcium carbonate or sulphate.
[Edited on 4-7-2015 by S.C. Wack]
|
|
|
glassplass
Harmless
Posts: 6
Registered: 16-3-2016
Member Is Offline
Mood: No Mood
|
|
| Quote: Originally posted by madscientist | | .....The first method was heating calcium sulfite and sodium nitrate together. This had seemingly good yields of sodium nitrite..... I prepared the
CaSO3 from NaHSO3 and CaCl2..... |
Can you use potassium metabisulfite or potassium sulfite instead of calcium sulfite?Probably you can but both nitrite and sulfite ar soluble in
water,and its dificult to separate them...
I dont have any sodium sulfites so i look for an alternative.
Can you directly buble SO2 gas into the solution of Ca(OH)2 or CaCO3,but aggain solubility in water is a problem so it will be ineffective,or?
If you cant do anything,than SO2 needs to be bubled into solution of NaOH to get mixtures of sodium hydrogen sulfite and sodium sulfite,but i dont
know when is enough,it will be contaminated with hydroxide.
Sorry for confusion,help!! 
|
|
|
AJKOER
Radically Dubious
Posts: 3026
Registered: 7-5-2011
Member Is Offline
Mood: No Mood
|
|
| Quote: Originally posted by AJKOER | .........
Here is another old source, "Journal of the Society of Chemical Industry", Volume 27, based on commercial processes with some very insightful
commentary on processes (see pages 484 to 485, link: http://books.google.com/books?pg=PA484&lpg=PA484&dq=... ). For example:
"Thermal decomposition.—The alkali nitrates when heated above their fusion point evolve oxygen and furnish nitrites, but this reaction is, of no
practical importance owing to the simultaneous occurrence of a further decomposition to oxide. "
There is also an interesting aqueous method mentioned. To quote:
"The heating of a concentrated solution of lead nitrate with finely divided lead leads to the formation of insoluble basic lead nitrite, which
furnishes sodium nitrite by double decomposition with sodium carbonate. This reaction is, however, only of theoretical interest. "
...... |
Apparently, aqueous nitrate can be reduced by Pb to nitrite:
Pb + NO3- → PbO + NO2- (see reference [22] below)
Also, nitrite then reduced further by Pb to N2 or NH3 (depending on pH).
2 NO2- + 3 Pb + H2O → N2 + 3 PbO + 2 OH-
NO2- + 3 Pb + 2 H2O → NH3 + 3 PbO + OH-
Reactions reference:: https://www.google.com/url?sa=t&source=web&rct=j&...
Cited reference[ 22] is Uchida, Miho; Okuwaki, Akitsugu. (1998), "Decomposition of nitrate by in situ buff abrasion of lead plate",
published in Hydrometallurgy, 49, 297-308. To quote the abstract:
"A new approach to the decomposition of nitrate ion using Pb metal has been developed. Removal of the oxide layer formed on the surface of Pb plate
and the production of Pb powder have been achieved by in situ abrasion of Pb plate in NH4NO3 solution with an abrasive buff. NO3− was reduced to
NO2− at initial concentrations of NH4NO3 in the range 0.01–0.2 M, rotational speed of buff, 100–800 rpm and temperature, 25–80°C. Complete
reduction of NO3− in 0.05–0.2 M NH4NO3 solution was achieved at 80°C within 4 h. As the temperature increases, the reduction rate of NO3− to
NO2− increases abruptly. The reduction rate increases gradually with rotational speed. The formation of NO2− is almost independent of the initial
NO3− concentration. Reduction of NO3− to NO2− is related to corrosion of Pb in NH4NO3 solution."
Link: http://www.sciencedirect.com/science/article/pii/S0304386X98...
So apparently buffing the Pb electrode to remove PbO avoids the formation of the basic lead nitrite and leaves aqueous nitrite.
So no extreme temperatures or lead vapors are required, just combine a buffing hand tool with a galvanic corrosion reaction (which, I suspect would be
accelerated by the addition of NaNO3 to the aqueous NH4NO3).
|
|
|
aga
Forum Drunkard
Posts: 7030
Registered: 25-3-2014
Member Is Offline
|
|
Preparation of sodium nitrite
To attempt to follow the Congo Red sysnthesis, 0.5g of sodium nitrite is required, hence this preparation.
According to this procedure :
http://www.prepchem.com/synthesis-of-sodium-nitrite/
71g of sodium nitrate and 172g of elemental lead was added to an iron vessel, then heated strongly with a blowtorch.
(this represents an excess of lead).
The mixture was stirred with an iron bar throughout the process.
After 20 minutes of heating & stirring, a pink scum was evident on the surface of the liquid, presumably PbO.
After 30 minutes the mixture began to take on the appearance and texture of pink blancmange, with many bubbles in the structure: a foam.
At 40 minutes the centre of the mass began to glow red, so heating was discontinued, although manual stirring was maintained.
After 1 minute the mass began to solidify, so stirring was continued to try to achieve a small fragment size as per the reference.
The cooled mass appeared as a yellow/ochre 'rocks' with some pink showing, mostly in areas furthest away from the flame.
The product was extracted with a total of 500ml of hot water, which had to boiled down to about 50ml before it would crystallise out.
Commercial test strips were used to test for the presence of Nitrites
The plastic cup used for weighing the NaNO3 had 100ml of DW added, and showed positive for Nitrates, negative for Nitrites.
1.5ml of the supernantant liquid from the interim product was added to the same plastic cup, then the test was repeated with this mixture and a very
strong Nitrite result was seen.
(The nitrite block is second from the left.)
The crude product was recrystallised and dried, giving a yield of 26.32g (44%)
|
|
|
Σldritch
Hazard to Others
Posts: 309
Registered: 22-3-2016
Member Is Offline
Mood: No Mood
|
|
Im trying too make nitrite by reduction with sulfur (2 NaNO3 + S + NaX --> 2 NaNO2 + NaSO4 + X2) and i have had some minor success with igniting
what is essentially yellow powder with extra nitrate but i am trying to do this in solution now because of the explosion hazard.
When i tried melting sodium hydroxide and sulfur together i got red to orange goop and i poured it out on a rusty iron plate to cool to later add it
to a solution of potassium nitrate. When i did i obtained a very dark green solution which mostly dispeared on boiling down the solution.
So what is the green compound? Iron complex with polysulfide or nitrite? Iron nitrosyl is brown and the ph is really high so i doubt that it would
exist in the solution. I will post progress if i make any.
Sorry if this is supposed too go in beginnings not sure tough.
|
|
|
Melgar
Anti-Spam Agent
Posts: 2004
Registered: 23-2-2010
Location: Connecticut
Member Is Offline
Mood: Estrified
|
|
| Quote: Originally posted by Σldritch | Im trying too make nitrite by reduction with sulfur (2 NaNO3 + S + NaX --> 2 NaNO2 + NaSO4 + X2) and i have had some minor success with igniting
what is essentially yellow powder with extra nitrate but i am trying to do this in solution now because of the explosion hazard.
When i tried melting sodium hydroxide and sulfur together i got red to orange goop and i poured it out on a rusty iron plate to cool to later add it
to a solution of potassium nitrate. When i did i obtained a very dark green solution which mostly dispeared on boiling down the solution.
So what is the green compound? Iron complex with polysulfide or nitrite? Iron nitrosyl is brown and the ph is really high so i doubt that it would
exist in the solution. I will post progress if i make any.
Sorry if this is supposed too go in beginnings not sure tough. |
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.
|
|
|
| Pages:
1
..
7
8
9
10
11
..
16 |
|
![[*]](images/xpblue/default_icon.gif){kind=icon}
