09-10-03 10:06
No 458114
After reading much about the preparation for the
aforementioned chemical, I'm still at a loss for
what seems to be a suitable method. I found a
pyrotechnic board discussing this very subject
(interestingly, they linked to the hive during the
course) and it was their opinion that using
sodium nitrate to oxidize lead in an iron crucible
is an inferior technique. I found one other,
that of passing nitric oxide through a NaOH solution,
but it requires large quantities of copper, which may
or may not be a problem. In researching the production
of nitric oxide (also pyro discussions) they mentioned
using starch in dilute (5M) nitric acid would also
generate it.
(I also found information about burning air with
excess oxygen at various temperatures to produce
this gas, but that's a bit outside of my capabilities
at the moment)
Could anybody with information regarding these two
methods [Nitric Acid + (starch || metallic copper)]
specifically the ratios of the NOx's produced provide
some insight? Also, if anybody has info on any other
procedure that might prove valuable, your input would
be greatly appreciated.
(Yes, I admit it, I suffered a moderate bout of cranial
flatulence and briefly (20min) confused the nitrite and
the sulfite, which led me on a rollercoaster of emotion)
over and under, then back to the start
this sidewalk's bottom is the same as its top
(Stranger)
09-10-03 20:18
No 458202
I remember that alkali metal nitrates lose oxygen to form nitrites if heated strongly: 2 KNO3 --> 2 KNO2 + O2 (I don't have the details, but as far as I remember it's like 500 or 600 degrees C or so.). I think the reaction is reversible, so it might have to be done under vacuum. Just some thoughts to follow up.
(Hive Addict)
09-10-03 20:29
No 458206
http://www.geocities.com/dritte123/NaNO2
(Hive Bee)
09-11-03 00:56
No 458249
(Rated as: excellent)
note: some of the equations are probably erroneous due to blurry text, I didn't have time to properly balance them out.
Sodium Nitrite, NaNO2
After removing the small dome from the inner file-clay mantle of a Rossler gas-furnace, place upon the mantle a strong iron-wire triangle, and set upon the triangle a shallow iron dish (2.5cm high, 12cm upper diameter) having a smooth bottom. Place 85 grams of Sodium Nitrate (NaNO3) in the dish and close the furnace. As soon as the dish has become faintly incandescent and the molten nitrate just begins to give off bubbles of oxygen, gradually add 206 grams of lead in the form of old pieces of sheet lead or lead tubing. The lead is at once vigorously oxidized, and, if sturred continually with an iron spatula, becomes almost completely converted into oxide of lead in half an hour. Empty the contents of the small iron dish into a large deep iron one, and repeat the operation several times, using the same amounts of Sodium Nitrate and lead. Place the various products in the large iron dish, extract once with boiling water, and decant upon a creased filter. Dry the reside of lead oxide and set aside for the experiment on page 52. Pass a strong current of CO2 into the still boiling-hot filtrate,for a few minutes only, filter off the lead carbonate which separates, and neutralize the solution while stirring it, by carefully adding nitric acid from a pipette or burette. Evaporate the solution to crystallization. The crystals which separate first, consist partly of nitrate and may be used again for remelting with lead; the mother-liquor gives pure nitrite. A normal solution of the nitrite is prepared by dissolving 69g of it in water and diluting to one liter.
Reaction: NaNO3 + Pb = NaNO2 + PbO.
A small part of the nitrite is converted into sodium plumbite, which must be decomposed by carbon dioxide - nitrous acid reacts with peroxides and with permanganate according to the following equations:
MnO2 + HNO2 + HNO3 = Mn(NO3)2 + H2O
2 HMnO4 + 5 HNO2 = 2 Mn(NO3)2 + HNO3 + 3 H20
with sulphanilic acid it gives diazobenzenesulphonic acid:
C5H4(NH2)SO3H + HNO2 = C8H4(N=N)(SO2) + 2 H2O
Test: Sodium Nitrite crystals, upon treatment with dilute mineral acids, go very easily into solution, with considerable effervescence and evolution of nitrous anhydride. The acid solution turns potassium iodide and starch paper blue and dissolves the peroxides of manganese and lead with great rapidity. To test the strength of the sodium nitrite, fill a burette with the normal solution and run it into a solution of 11.55 grams of the sodium salt of sulphanilic acid, C6H4NH2.SO2Na + 2 H2O which is cooled with ice and made strongly acid with hydrochloric acid; stir constantly, and continue the addition of the nitrite solution until a drop of the liquid gives a strong blue color with potassium iodide and starch paper. If the nitrite is pure, 50 mL are necessary. Or, 50 mL of the nitrite solution may be measured into a graduated flask by means of pipette, diluted to one liter, and this dilute solution run from a burette into tenth normal permanganate with is made acid with sulphuric acid and warmed to 40-50°C, stirring during the addition. Just 50cc of the twentieth normal nitrite solution should be used for 50cc of the permanganate.
Detection of Nitrous Acid -
(i) On adding dilute sulphuric acid, potassium iodide, and a solution of starch to a nitrite, a blue colour develops, a liberation of iodine which combines with the starch. Nitrates produce a similar color only after standing for some time.
(ii) Ferrous sulphate and acetic acid produce a brown color which is discharged on warming, brown fumes being evolved (compare detection of nitric acid, p.453)
(iii) Meta-phenylenediamine in acid solutions gives a brown colour with nitrous acid.
Estimation of Nitrous Acid and Nitrites -
(i) Nitrites can be estimated by a reduction to nitric oxide as in the case of nitric acid.
(ii) When present in smaller quantities, the nitrite may be reduced to ammonion estimated by Nessler's solution or by titration as in the case of Nitric Acid.
(iii) In the absence of other oxidisable compounds, nitrous acid (e.g.) in "waste acids" from nitration) can be estimated by oxidation with potassium permanganate. 2 KMnO4 = 3 H2SO4 + 5 HNO3 = K2SO4 + 2 MnSO4 + 5 HNO3 + 3 H2O
(iv) Traces of nitrites may be estimated by means of the colour produced with meta-phenylenediamine.
over and under, then back to the start
this sidewalk's bottom is the same as its top
(Hive Bee)
09-12-03 07:49
No 458553
I refer to: ../rhodium /nitrous
Is there any way to optimize this reaction for the
formation of nitric oxide?
over and under, then back to the start
this sidewalk's bottom is the same as its top
(Chief Bee)
09-12-03 09:25
No 458574
Nitric oxide is NO - the synthesis at my page only produces nitrous oxide (N2O).
(Hive Bee)
09-12-03 10:08
No 458584
It also mentions that if the reagents are added in specific
orders, varying amounts of nitric oxide and nitrogen dioxide
are formed, which suggests to me that it might be possible
to optimize it for one or the other... I could be horribly
wrong, however....
over and under, then back to the start
this sidewalk's bottom is the same as its top
(Hive Bee)
09-13-03 05:31
No 458789
gleaned from http://www.ucc.ie/ucc/depts/chem/dolchem
Nitric Oxide, NO is prepared by the action of Copper, Cu, or Mercury, Hg, on dilute Nitric Acid, HNO3, and was called Nitrous Air.
3Cu + 8HNO3 ==> 3Cu(NO3)2 + 2NO + 4H2O
from http://www.finishers-management.com/may2
The most common method for controlling NOX emissions is gas absorption scrubbing with sodium hydroxide (NaOH) as the scrubbing medium. NO2 gas possesses rather high solubility and reactivity in aqueous or alkali solutions such as NaOH. Gaseous NO, on the other hand, is only slightly soluble in water and not very reactive in aqueous or alkali solutions. In addition to the solubility problems with NO, when an NO/NO2 waste gas stream is brought into contact with an alkali solution, nitrates and nitrites can be formed, further hindering the absorption of the NOX. Further complicating the absorption in an alkali solution is the fact that the sodium nitrite formed can further react with the remaining NO2.
from http://www.ucc.ie/ucc/depts/chem/dolchem
Similarly, nitric oxide reacts with sodium hydroxide in aqueous solution to produce a mixture of sodium nitrate, sodium nitrite and Water.
2NO2 + 2NaOH ==> NaNO2 + NaNO3 + H2O
from http://www.uq.edu.au/_School_Science_Les
13.11.1 Make nitrogen monoxide
See diagram 13.11.1 | See 12.3.11: Reaction of dilute nitric acid with metals
Nitrogen monoxide reacts with oxygen existing in air so easily and rapidly to produce a reddish brown gas, nitrogen dioxide, that the formation of colourless nitrogen monoxide could hardly be perceived directly through experimental phenomena. To solve this difficulty, an apparatus is designed as follows:
Fit two ends of a large glass tube, that has a diameter of 1.5-2 cm and a length of about 15 cm, with two 1-hole stoppers respectively carrying a small glass tube.
Connect one of the two small glass tubes to a short rubber tube with a pinch cock on it.
Place a rubber gasket of 2-3 mm thickness with several small holes drilled in it in the middle part of the large glass tube.
Drop a few pieces of copper on the rubber gasket.
Take the stopper off from the other end of the large tube and add diluted nitric acid of 1:3 concentrated nitric acid : water.
Make enough to make the solution level near the mouth of the glass tube, not leaving any air bubbles in the tube after replacing the stopper..
Turn the tube upside down, clamp it vertically on an iron stand, and place a small beaker under the tube to receive the solution coming out from the tube.
During the reaction the colourless gas produced gradually presses the solution out of the tube.
When the copper pieces are no longer in contact with the nitric acid solution, the reaction stops, leaving colourless nitrogen monoxide collected in the upper portion of the glass tube and a blue solution in the small beaker.
After taking the pinch cock off the rubber tube to let air enter the tube, the nitrogen monoxide inside the tube can be seen to be oxidized.
A slowly rising level of the solution shows that nitrogen dioxide is liable to dissolve in water.
In order to prevent air contamination, absorb the nitrogen dioxide in the tube with alkali solution.
Could anybody draw a diagram of that last one?
Maybe my brain isn't working right, but it seems like
the HNO3 would start pouring through the gasket as
soon as you put it in the tube....
over and under, then back to the start
this sidewalk's bottom is the same as its top
(Hive Bee)
09-14-03 06:41
No 458964
on the hive somewhere, that a simple route to NaNO2 via NaNO3 was to melt the NaNO3- mp about 300C. bubbles of O would then form. the heating of the liquid is then continued until no more bubbles of O evolve, and you have NaNO2. does that sound unrealistic?
"Im not as think as you stoned I am"
(Hive Bee)
09-14-03 08:14
No 458972
Preparation of sodium nitrite (http://www.sciencemadness.org/talk/view
This site rocks... These people are even more insane
than us bees... Great inorganic synths...
over and under, then back to the start
this sidewalk's bottom is the same as its top
(Newbee)
09-18-03 13:05
No 459743
A method which seems interesting to me is passing dry SO2 over a fused mixture of NaNO3 and CaO - SO2 reduces the nitrate to nitrite, and the SO3 imediately combines with the CaO to form CaSO4, which is practically insoluble in water. So you can just wash out the nitrite and evap to form nice pure crystals. The SO2 can be formed in an easily controlable reaction by the action of acid on sodium metabisulphite (home-brewing suppliers...) and dried with H2SO4.
Using KNO3 might be better than NaNO3, since there is more of a difference in solubility between KNO2 and KNO3, making seperation by fractional crystalisation a little easier.
KNO2 can also be made by mixing KNO3 and C, roughly 17:2, and igniting it. The C acts as a reducer, and the heat also acts to decompose the nitrate. Use it in a fairly thin layer in a large, covered vessel (clay plant pot or something similar). Seems simple enough, even if yields aren't great. I think I read about this one on sciencemadness.org, or it may have been roguesci.org...
Heating NaNO3 or KNO3 with CuSO4 or FeSO4 or other transition metal salt (although those are the ones that I have seen mentioned) forms the TM nitrate in situ, and these are easily decomposed by heat, forming NO2, which can then be bubbled through KOH soln to form the nitrate and nitrite, which can be seperated by fractional crystalisation. This is very similar to how nitric acid was made by alchemists etc hundreds of years ago. I prefer this method of NOx production to carbohydrate/HNO3 or metal/HNO3 because it doesn't waste nitric acid.
(Hive Bee)
09-23-03 13:03
No 460533
Kirk-Othmer:
Manufacturing
Sodium nitrite has been synthesized by a number of chemical reactions involving the reduction of sodium nitrate [7631-99-4], NaNO3. These include exposure to heat, light, and ionizing radiation (2), addition of lead metal to fused sodium nitrate at 400-450°C (2), reaction of the nitrate in the presence of sodium ferrate and nitric oxide at (2), contacting molten sodium nitrate with hydrogen (7), and electrolytic reduction of sodium nitrate in a cell having a cation-exchange membrane, rhodium-plated titanium anode, and lead cathode (8).
Industrial production of sodium nitrite is by absorption of nitrogen oxides (NOx) into aqueous sodium carbonate or sodium hydroxide. NOx gases originate from catalytic air oxidation of anhydrous ammonia, a practice common to nitric acid plants:
Gas contact is typically carried out in absorption towers over which the alkaline solutions are recirculated. Strict control over the conditions of absorption are required to efficiently capture the NOx and convert it predominantly to sodium nitrite according to the following reaction, thereby minimizing the formation of by-product sodium nitrate. Excessive amounts of nitrate can impede the separation of pure sodium nitrite from the process.
Solutions of sodium nitrite thus produced are concentrated and a slurry of crystals obtained in conventional evaporation and crystallization equipment. Much of this equipment can be of mild steel construction because sodium nitrite functions as a corrosion inhibitor toward most ferrous metals. The crystals are typically separated from the mother liquor by centrifugation and subsequently dried. Because of its tendency to lump and cake rapidly in storage, dry sodium nitrite products are frequently treated with an anticaking agent to keep them free-flowing. Alternatively, larger flakes or pellets are prepared from the granular material through a compaction process. The limited surface contact between these larger particles allows them to remain uncaked for extended periods. Technical solutions for commerce can be obtained directly from the process; higher purity solution products are prepared by dissolving crystals.
[2] J. W. Mellor, A Comprehensive Treatise on Inorganic and Theoretical Chemistry, Vol. 8, Longmans, Green & Co., London, 1928; J. W. Mellor, Supplement to Mellor's Treatise on Inorganic and Theoretical Chemistry, Vol. VIII, Suppl. II, Part II, John Wiley & Sons, Inc., New York, 1967.
[7.] Patent US2294374 (Sept. 1, 1945), J. R. Bates (to Houdry Process Corp.).
[8.] Patent DE2940186 (Apr. 24, 1980), M. Yoshida (to Asahi Chemical Industry Co., Inc.). = Patent US4312722
(Hive Bee)
09-27-03 20:33
No 461360
kingspaz from madscience board quoted as saying:
'i make mine like this:
2Al + 3KNO3 ---> Al2O3 + 3KNO2
it needs to be molten about 40 minutes and its done. i think its a decent method just for the ease of doing it. my KNO2 is a very pale yellow/white.'
'xoo, my method above seems to work quite well. the trick is to keep the KNO3 just above its melting point (334*C if i remember correctly). this can of course be judged by the KNO3 melting'
Is there any benefit extending the rxn time?
Will rxn time or excessive heat cause an explosion? or both?
(Hive Bee)
10-10-03 12:00
No 463761
Abstact:
Nitrates, such as alkali metal nitrates and ammonium nitrate, are electrolytically reduced to the corresponding nitrites by a process wherein an aqueous solution containing a nitrate is supplied into a cathode chamber of an electrolytic cell including cathode and anode chambers separated by a cation exchange membrane and an electric current is applied to the electrolytic cell, while maintaining the pH of the aqueous solution at a value of at least about 4. This process can be advantageously applied not only to the manufacture of nitrites but, also, to the treatment of waste nitrates. In the treatment of waste ammonium nitrate, the ammonium nitrite so formed is conveniently further treated by subjecting the electrolytically reduced catholyte to thermal decomposition outside the electrolytic cell, and removing the so formed nitrogen and water from the reaction system.
roger2003
(Stranger)
10-10-03 23:42
No 463846
does anybody know the standard potential of sodium nitrite?
i was wondering if a partially oxidized metal salt (e.g. Fe(+2) )could be used to reduce NaNO3 in an aqueous soltion.
with the standard potential i could see which anions would possibly work.
(Hive Bee)
11-17-03 00:25
No 471276
(Rated as: good read)
Nitrites from Nitrates: Patent #792,515 & Wild Conjecturing
First the patent. Seen on another thread in the hive, acquired and typed for all to enjoy.
Jacob Grossmann, of Manchester, England.
Process of making nitrites.
Be it known that I, Jacob Grossman, a subject of the King of England, residing at Manchester, in the county of Lancaster, England, have invented a new and useful Improvement in the Manufacture of Nitrites, of which the following is a full specification (yeah, right).
It is known that when nitrate of soda is fused with carbon in the presence of caustic alkali, nitrite of soda is produced more or less according to the following equation:
2 NaNO3 + 2 NaOH + C ----> 2 NaNO2 + Na2CO3 + H2O
As caustic soda has to be used in excess, the resulting melt contains four salts which are easily soluble and which on lixiviation go into solution -- vis., sodium nitrite, nitrate, carbonate, and hydrate -- besides silicates, and other impurities. This interferes seriously with the subsequent separation and purification of the nitrite. In order to overcome this objection, I use the oxid or hydrate of an alkaline earth instead of caustic alkali. The reaction may be represented by the following equation:
2 NaNO2 + Ca(OH)2 + C ----> 2 NaNO2 + CaCO3 + H2O
It will be seen that the only possible substances which on dissolving the melt can go into solution are the nitrite formed and whatever nitrate has not been decomposed, and as the resulting liquor may be boiled down to dryness and melted again with more caustic earth and carbon, the ultimate products obtained in solution will be nitrite of soda and a small percentage of undecomposed niter free from other soluble compounds. Instead of the oxid or hydrate of calcium the oxid or hydrate of barium, strontium, or magnesium may be used. A small proportion of caustic alkali may be mixed with the lime or added to the melt to act as a carrier. The yield of nitrite and the loss in niter by overreduction and similar causes depends on the allotropic form of carbon used. The yield is highest and the loss least in the case of graphite, natural or refined or artificial. Coke comes next and other forms of carbon after that.
As an example, I may describe how this process may be worked in the case where lime and graphite are used. A quantity of niter by itself or mixed with lime is melted in an iron pot provided with a mechanical stirrer and a mixture of graphite and slaked lime (to which may be added a little caustic soda) added gradually until the mass assumes a yellow color and shows by test that no more decomposition to nitrite takes place. The mass after cooling is lixiviated and the liquor separated from the insoluble residue. The liquor is boiled down to dryness in an open or in a vacuum pan, and the dry mass transferred to the melting-pot and fused and treated again with a further quantity of lime and graphite. The melt resulting from this may be again lixiviated, boiled down, and melted with lime and graphite. From the resulting solution, which contains practically only nitrite and a small quantity of nitrate, commercially-pure nitrite can be easily prepared in a manner well known to chemists. The reaction does not take place exactly in the proportions shown in the equation. More lime and carbon have to be used than would correspond to the quantity of nitrite produced. This process may be reversed in this way that the nitrates of the alkaline earths may be fused with caustic alkali and carbon. Inert substances, such as chlorid of sodium, may be added as diluents. Nitrate of potash may be used instead of nitrate of soda.
Having now particularly described and ascertained the nature of my said invention and in what manner the same is to be performed, I declare that what I claim is--
1. In a process for the manufacture of alkaline nitrites by reducing nitrates with carbon, the fusing of the nitrates with graphite in the presence of an oxid of an alkaline earth, substantially as described.
2. In a process for the manufacture of alkaline nitrites by reducing nitrates with carbon, the fusing of the nitrates with graphite in the presence of a hydrate of an alkaline earth, substantially as described.
3. In a process for the manufacture of alkaline nitrites by reducing nitrates with carbon, the fusing of the nitrates with graphite in the presence of an oxid and a hydrate of an alkaline earth, substantially as described.
In witness whereof I have hereunto set my hand in presence of two witnesses.
Jacob Grossman.
Next, some useful numbers, and conjecturing:
MW MP Note
NaNO3 85 308
NaNO2 69 271 dec. from 320!
NaOH 40 318
Na2CO3 106 851 dec. from 400
KNO3 101 333
KNO2 85 441 dec. from 350
KOH 56 360
K2CO3 138 891
Ca(OH)2 74 dec. 580 --> CaO + H2O
CaCO3 100 dec. 825 --> CaO + CO2
If this process has been attempted before by anybee, (seem to remember it having been tried), unless the temperature was controlled very carefully, they would likely lose much of their nitrite to decomposition. So let's put that table into a different form:
NO2 dec. NO2 melts Margin, NO2 NO3 melts Margin, NO3-NO2
Na: 320 271 49 308 12
K: 350 441 -91 333 17
12 degrees of margin between melting and decomposition!
17 degrees of margin between melting and decomposition!
Bees, you better lower the heat as soon as that sucker melts, else you're gonna lose some nitrite! And don't do this without some ventilation! Ning doesn't know what NO2 will decompose into, but suspects it goes something like 2 NaNO2 --> Na2O + NO + NO2 // N2O3ˇ¦Not friendly!
So, if you use NaNO3, you have less margin between melting of NO3 and decomposition of NO2, but the NO2 will be melted also, which should make the melt work better.
If you use KNO3, you have 5 degrees more margin, but the NO2 formed will not be molten, meaning as the reaction progresses, it will get more and more stiff and dry. This might lower yields. It would be a good method of monitoring the reaction progress, howeverˇ¦
With these caveats and dangers in mind, seems like an awfully OTC method to nitrites. Experimenter bees, have at it!
ps. Why in fuck's name is the "pre" tag not working?
(Hive Addict)
11-20-03 18:12
No 472000
I have been looking into various methods outside of the NaNO3/Pb method for preparing sodium nitrite. The first method was heating calcium sulfite and sodium nitrate together. This had seemingly good yields of sodium nitrite; however, most of the sodium nitrite was destroyed when I forgot that I was boiling off the water on an electric burner outside, and it was fried for around 10 minutes... I prepared the CaSO3 from NaHSO3 and CaCl2. The procedure for preparing CaSO3 was mixing stoichemical amounts of NaHSO3 and CaCl2, then wetting the mix; SO2 gas was liberated, CaSO3 was formed, as well as NaCl. The sludge left over was scraped into a filter, then was run through the sludge, which removed the sodium chloride. Of course, SO2 gas, generated by reaction of the NaHSO3 with an acid, could have been bubbled through a solution of Ca(OH)2 to form CaSO3 as well.
This is what I tried:
I placed 15g of CaCl2 in a beaker, along with 28.1g NaHSO3; then added 50mL of hot water. It fizzled, and liberated a good-sized amount of SO2 gas. A few minutes after SO2 gas was no longer being liberated, I dumped the contents of the beaker into a filter. I poured an additional 100mL of water through the filter, to insure that very little NaCl was left. I lost some CaSO3, due to CaSO3 being somewhat soluble in acids (and the solution formed after reaction of the CaCl2 and then NaHSO3 was obviously sulfurous acid, due to dissolved SO2 gas). I then heated the contents of the beaker with a "hotplate" that warms to about 120C, to dry it. I soon had a yellowish sludge, which is Ca(HSO3)2, which is simply CaSO3 and H2SO3 bonded together. After a while, all of the sulfurous acid was finally driven off, and I was left with dry CaSO3. I had 13.3g of CaSO3; an 82% yield. I then proceeded to heat the 13.3g of CaSO3 with 9.4g of NaNO3, on a propane burner. I stirred constantly to insure even heating. After about five minutes, I had a deep-yellow solid mix; the color didn't change any more after that point. I dumped the contents of the beaker into 500mL of water, and filtered. I began heating the filtered solution (which was a golden yellow), but I forgot about it... and ended up toasting the NaNO2 for quite a while with that electric burner. I'm confident it was destroyed by that, because when I try to dissolve that solid in water, I just get a brownish mix; the color of the solid reminds me of soil high in clay content.
This is my second idea for preparing sodium nitrite; this one has not been tested as of yet, but will be shortly. It requires sulfamic acid (HSO3NH2), calcium oxide, and sodium nitrate.
CaO + 2HSO3NH2 --> Ca(SO3NH2)2 + H2O
This reaction will have to use alcohol as the solvent, not water. Sulfamic acid hydrolizes to ammonium hydrogen sulfate; calcium sulfamate hydrolizes to calcium ammonium sulfate...
Ca(SO3NH2)2 + CaO + 3NaNO3 --> 2CaSO4 + 3NaNO2 + N2 + 2H2O
Ca(SO3NH2)2 is very water soluble. CaSO4 is not. Since the calcium sulfate formed in the second reaction would not dissolve in water, while the sodium nitrite would, extracting the sodium nitrite would be very easy.
made some sodium nitrite today. This is the process I used:
2HSO3NH2 + 2CaO + 3NaNO3 ----> 2CaSO4 + 3NaNO2 + 3H2O + N2
I thoroughly mixed 20g of powdered NaNO3, 8.8g of powdered CaO, and 15.2g of powdered HSO3NH2 (sulfamic acid). I then began heating it on a propane burner, stirring rapidly to insure even heating. After a few minutes, the mixture suddenly began fizzling and billowing water vapor. I stopped heating the beaker, and the reaction continued to accelerate. After about three minutes, it stopped reacting. The reaction certainly was more exothermic than I expected. I suppose a stoichemitric amount of Ca(OH)2, or even CaCO3 could be substituted for the CaO - but the reaction would be much less exothermic.
After the contents of the beaker cooled down, I then poured 200mL of hot water into it. It fizzled quietly (I'm hypothesizing that was a small amount of leftover sulfamic acid hydrating, forming the ammonium ion which was being oxidized to nitrogen gas by the nitrite ion), but soon stopped. I poured it through a filter. A considerable amount of CaSO4 was caught in the filter. The filtered solution was a beautiful light, golden yellow. I'm currently heating it gently in a flask (don't want the nitrite ion being oxidized by oxygen gas), at about 90C, to drive offThe theoretical yield was about 16.2g of NaNO2. I managed to scrape about 11g of NaNO2 off of the bottom of the flask. The NaNO2 is a very pale yellow.
water. Report on the yield is on the way.
The theoretical yield was about 16.2g of NaNO2. I managed to scrape about 11g of NaNO2 off of the bottom of the flask. The NaNO2 is a very pale yellow.
Here's an idea that occurred to me a while back for preparing nitrites (this should also work for preparing nitrous acid, of course).
2NaHSO3 + 2NaNO3 --(heat)--> 2NaHSO4 + 2NaNO2 ----> 2Na2SO4 + H2O + NO2 + NO
The sodium bisulfate won't be able to liberate a significant quantity of nitric acid, as the pKa is too high (not acidic enough). The generated gas should be almost entirely composed of equamolar quantities of NO2 and NO.
Basically the idea is to mix equamolar quantities of sodium bisulfite powder and sodium nitrate powder, heat in a flask, and bubble the generated gasses into a strongly alkaline solution, yielding relatively pure nitrite (or into water to yield metallic cation-free nitrous acid). The generated gasses must not be allowed to come into contact with atmospheric oxygen, or much of the nitric oxide will be oxidized to nitrogen dioxide, rendering the gasses nearly useless for preparing relatively pure nitrite (unless one plans on attempting fractional crystallization).
Edit: The edit button works!
Ramiel, this is my hypothesis as to what the reaction was between potassium nitrite, aluminum, and (I assume) water that generated ammonia:
2KNO2 + 4Al + 6H2O ----> 2KOH + 2Al2O3 + 2H2O + 2NH3
Some more ideas:
Na2S2O4 + 3NaNO3 ----> 2Na2SO4 + NaNO2 + NO + NO2
Na2S2O4 is the ingredient in many solid toilet bowl cleaners (known as "sodium hydrosulfite.")
3NaNO3 + S ----> Na2SO4 + NaNO2 + NO + NO2
Those reactions could be of interest if they proceed as predicted due to the possiblity of providing a dry equamolar mixture of NO and NO2. Isolating the solid nitrite produced surely would prove to be difficult.