http://www.geocities.com/dritte123/NaNO2.html (http://www.geocities.com/dritte123/NaNO2.html)
I refer to: https://www.thevespiary.org/rhodium/Rhodium/chemistry/nitrous.html (https://www.thevespiary.org/rhodium/Rhodium/chemistry/nitrous.html)
Is there any way to optimize this reaction for the
formation of nitric oxide?
gleaned from http://www.ucc.ie/ucc/depts/chem/dolchem/html/comp/nitric.html (http://www.ucc.ie/ucc/depts/chem/dolchem/html/comp/nitric.html)
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/may2002/nox.htm (http://www.finishers-management.com/may2002/nox.htm)
:
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/html/dict/000a2.html (http://www.ucc.ie/ucc/depts/chem/dolchem/html/dict/000a2.html)
:
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_Lessons/topic13.html#13.11.0 (http://www.uq.edu.au/_School_Science_Lessons/topic13.html#13.11.0)
: (Also has many good gas generating procs)
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....
Preparation of sodium nitrite (http://www.sciencemadness.org/talk/viewthread.php?tid=52)
(http://www.sciencemadness.org/talk/viewthread.php?tid=52)
This site rocks... These people are even more insane
than us bees... Great inorganic synths...
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 (http://l2.espacenet.com/dips/viewer?PN=US2294374&CY=gb&LG=en&DB=EPD)
(Sept. 1, 1945), J. R. Bates (to Houdry Process Corp.).
[8.] Patent DE2940186 (http://l2.espacenet.com/dips/viewer?PN=DE2940186&CY=gb&LG=en&DB=EPD)
(Apr. 24, 1980), M. Yoshida (to Asahi Chemical Industry Co., Inc.). = Patent US4312722 (http://l2.espacenet.com/dips/viewer?PN=US4312722&CY=gb&LG=en&DB=EPD)
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 // N2O3¡¦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, however¡¦
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?