Sciencemadness Discussion Board

N2O4 cheap and easy without H2SO4

franklyn - 21-10-2009 at 09:55

- About time there was a thread on Dinitrogen Tetroxide -

Dry distillation , why didn't anyone think of this before it's so stupidly simple.
Disclaimer : I have not actually done this , if in doubt wait for feedback ,
as they say , run it up the flag pole and watch who salutes it.

Ingredients - Fertilizer grades MnSO4•H20 ( Manganese not magnesium ) & NaNO3

*Note that crystallizing the salts from solution as described first is only necessary
if extremely pure analytic grade reagent Dinitrogen Tetroxide is wanted. Providing
you don't also want purified NaNO3 and MnSO4 for other uses , skip to part ( 2 )

- Solubility of NaNO3 in 100 grams of water increases from 40 grams at 0 ºC ,
to 170 grams at 100 ºC which is perfectly adequate. Adding more just raises
the boiling point so greater quantity can be dissolved at a higher temperature ,
at 200 ºC more than 80 % of the solution will be salt.

- Solubility of MnSO4 increases from 0 to 55 ºC decreasing at higher temperature.

( 1 ) Initial processing involves dissolving each salt separately into water at 100 ºC
then filtering each solution while hot to remove insoluble impurities. Next boil away the
water until it begins to show solid undissolved salt then let cool to precipitate some of
the crystals which are removed by filtering. More of the remaining water of the solution
is once again boiled away and cooled again to precipitate another batch of purified
crystals which are also removed by filtering. This routine is repeated for a third last
time to recover most all of the pure salt. Contaminants remain in solution as they
never reach saturation being only present in low concentrations.
Video demonstration and how to , here _

( 2 ) Dry NaNO3 by melting in an open pan in the kitchen oven at high heat ~ 315 ºC
for perhaps 20 minutes * carefully , above 380 ºC NaNO3 can decompose. If you can
perform this under vacuum it will be dry as you can get it. Moisture in store bought
appears as prills -
or as crystals -
The anhydrous ( NaNO3 ) is white powder.
This information indicates NaNO3 is safely melted without decomposition _

Dry MnSO4 by roasting in an open pan in the kitchen oven for about an hour at
high heat ~ 315 ºC ( anhydrous MnSO4 will not melt until it's at 710 ºC ).
If you can perform this with a vacuum desiccator it will be dry as you can get it.
Monohydrate ( MnSO4•H20 ) crystals are red colored , Tetrahydrate ( MnSO4•4H20 )
crystals are pink ( what you get after the previous crystalization proceedure )
The anhydrous ( MnSO4 ) is white powder.

Place each powder in a separate heavy plastic bag , the one it was bought in is good ,
then pound each into powder. Place into one bag , 155 grams of MnSO4 and add to
this 170 grams of NaNO3 and mix the two thoroughly by rolling it around for awhile.
( 151 grams MnSO4 is stoichiometeric , slight excess assures no NaNO3 decomposes )

( 3 ) This process is for dry distillation and requires high temperature resistent reaction
still. Pour the powder into the vessel in which it will be heated and run the vapor evolved
through a condensor first or expediently just into the closed receiver in an ice bath.
CAUTION : Be absolutely sure that any escape of NO2 will vent outdoors !

The reaction is :

2 NaNO3 + MnSO4 -> Na2SO4 + Mn(NO3)2 -> ( Mn(NO3)2 -> MnO2 + 2 NO2 )

2 NO2 on chilling condenses quantitatively 92 grams of N2O4

NaNO3 Initally melts above 308 ºC , but less than it's 380 ºC decomposition temperature.
The melt metathetcally becomes solid Na2SO4 with Mn(NO3)2 and that decomposes above
400 ºC into MnO2 + 2 NO2 . Here's the secret , MnO2 sequesters enough oxygen to avoid
N2O3 formation. MnO2 can only decompose above 535 ºC producing just 2 MnO and O2
The Na2SO4 won't even melt until 884 ºC , boiling at 1429 ºC
It's foolproof , there is no possibility of contaminating the anhydrous N2O4 .

N2O4 MSDS - M U S T . R E A D

Except for oleum a very similar notion by ' 497 '

Dry distillation notion for fuming Nitric acid by ' garage chemist ' here _
Call to action

Easy N2O4 can yield easy Ac2O , CL-20 anyone ? :)


Attachment: Manganese salts.pdf (314kB)
This file has been downloaded 1535 times

497 - 21-10-2009 at 17:54

I have produced quite large quantities of (somewhat impure) N2O4 by decomposition of the calcium nitrate + ammonium nitrate double salt that is available as a fertilizer in my location (and the rest of the US AFAIK). It consists of 5Ca(NO3)2 + 1NH4NO3 + 10H2O. On heating it melts in its own water of hydration, then boils off the water, turning to a paste.. It seems to boil off the water of hydration with very little decomposition of the nitrates. Then after heating to a higher temperature, the NH4NO3 decomposes releasing more water and a fairly small amount of NOx. Then on further heating it melts again and begins decomposing the Ca(NO3)2, forming large amounts of N2O4. The melt it not particularly corrosive, and can be contained in mild steel (although after repeated runs the container will eventually fail.) Stainless steel seems to hold up better.

I have done several batches using a large SS pot (with SS tubing welded to the top to make a sort of retort) containing around 5kg of the "CAN" fertilizer. It was heated on a homemade coal fired burner. Using a propane burner could not decompose it very quickly at all, and used huge amounts of propane. The N2O4 production proceeds fairly smoothly for several hours when using the coal burner.

The N2O4 production tapers off fairly quickly when completed. Interestingly the material left in the pot is not a white CaO powder as I expected, but a bright lime green liquid that solidifies to a brittle semi-translucent solid (still lime green) on cooling. This residue is very temperature resistant, it just sits there even when the bottom of the SS pot is red hot. I suspect it may be colored because of a small amount of metal ions dissolved off the walls of the pot, but I'm not sure. It is hygroscopic and becomes wet and pasty upon sitting out in the air for a while. It is partially soluble in water, but leaves a white insoluble residue. I am very curious to what the composition of this residue is, but I haven't had time to do any further analysis of it. I can speculate that it is some sort of basic nitrate that is more temperature stable.. except AFAIK no such compound exists for Ca (I know Mg does form basic nitrates, but never seen any for Ca.) And according to various publications Ca(NO3)2 should completely decompose.. So any ideas on that are welcome..

I think there is some water produced during the early part of the N2O4 production phase, but it does not seem to be much (I have yet to do a detailed analysis of the purity of the N2O4.) I did try condensing the N2O4 various ways. It is a challenge because of the O2 present. Gasoline cooled by a glycol-ice bath did condense some, but much was lost. I eventually tried bubbling it into solid CO2 slush in gasoline. That seemed to work well, although there were some issues with the bubbler tubing plugging up. More experiments are needed, but I haven't had time. I have yet to quantify the N2O4 produced, or the weight lost by the fertilizer.

That said, I do like your MnSO4 idea. If it does run smoothly and easily, making 100% N2O4, it will be even more useful. As long as it scales up well.. I would love to not need solid CO2 to condense the stuff! I'm afraid that the high fraction of solids in the melt may cause some problems though... it would definitely be nice to try it and find out. Unfortunately I'll probably have to order my MnSO4..

Also thanks for bringing up the N2O4 route to acetic anhydride, I missed that one. I'd love to try that some time. Ac2O makes such a lovely nitrating media..

watson.fawkes - 21-10-2009 at 19:30

Quote: Originally posted by 497  
This residue is very temperature resistant, it just sits there even when the bottom of the SS pot is red hot. I suspect it may be colored because of a small amount of metal ions dissolved off the walls of the pot, but I'm not sure. It is hygroscopic and becomes wet and pasty upon sitting out in the air for a while.
Iron is indeed the coloring agent in green glass. Most raw silica sources have iron contamination, and raw bottle glass is green unless decolorized. Calcium oxide is quite refractory, being the "lime" in "limelight", an old open-atmosphere incandescent light. It really seems like you've got mostly quicklime (CaO) left.

JohnWW - 21-10-2009 at 19:41

Amber glass, used mainly for beer and medicine bottles, is produced by adding ferric oxide (rouge) to molten glass. It is possible that ferrous oxide could produce a green color in glass, although I would think that something like chromium (III) or Cu(II) oxide would be better. Although many raw silica sand sources contain iron as an impurity, it is in the form of ferromagnetic minerals, particularly magnetite, produced like grains of silica as the impervious residue from the weathering of rocks, that can easily be removed by applying a strong magnetic field. The other main impurity in silica sand used for glassmaking would be carbonate minerals, such as calcite from sea-shells etc., but these are usually considered benign unless colored, and could be removed by dissolution with an acid anyway.

kilowatt - 21-10-2009 at 20:37

Would this reaction work with CaSO4 as well as MgSO4? Needless to say, this would be much cheaper.

Sedit - 21-10-2009 at 21:05

Where do you see MgSO4 anyware in that synthesis? Look again for safe measures.

497 - 21-10-2009 at 21:43

I agree, the green is probably from iron and/or chromium contamination.

It really seems like you've got mostly quicklime (CaO) left.

I don't think so.. CaO doesn't melt until 2572*C... This stuff was liquid at less than 900*C...

Would this reaction work with CaSO4 as well as MgSO4? Needless to say, this would be much cheaper.

It's MnSO4 not MgSO4.. Although you could try Ca and/or Mg, and it may still work, but you would get N2O4 contaminated with O2 (just like from pyrolysis of Ca(NO3)2) making it far harder to condense.. Mn has the useful ability to go from II+ to IV+ and in the process soak up that extra O by forming MnO2.. AFAIK no other metal will do that.. or at least no readily available one.

I must say franklyn, it is an ingenious idea. As long as molten NaNO3 will in fact react with MnSO4...

[Edited on 22-10-2009 by 497]

watson.fawkes - 22-10-2009 at 05:33

Quote: Originally posted by 497  

I don't think so.. CaO doesn't melt until 2572*C... This stuff was liquid at less than 900*C...
Mostly CaO, not entirely CaO. Fertilizers are generally crushed, prepared and/or treated ore. There's no economic point in refining the raw materials much. So I'd guess there's a few percent sodium and/or potassium in there, enough to partially flux the lime to cause some cohesion under heat, but not enough to prevent it from falling apart later.

franklyn - 23-10-2009 at 05:11

Quote: Originally posted by kilowatt  
Would this reaction work with CaSO4 as well as MgSO4? Needless to say, this would be much cheaper.

I know this is the third time it's mentioned in this thread , I made the same mistake
when doing research which is why I pointed out the difference at the start.
" Ingredients - Fertilizer grades MnSO4•H20 ( Manganese not magnesium ) & NaNO3 "

Since you mentioned it , MgSO4 goes for around $ 3 for 5 lb. in most any drug store.
If you believe this to be expensive perhaps you should take up a low overhead hobby
such as basketball.

Now consider MnSO4
5 pounds is 2268 grams , after removing water of hydration from MnSO4 there
remains 2018 grams which divided by the 155 weight for reaction is 13 portions
to 26 mols of NaNO3 which is 2210 grams or just 97.5 % of 5 ibs.
These quantities will produce 1196 grams N2O4 for under $ 11 worth
of materials by the lowest prices listed below ( if not obtained locally not
considering the cost of shipping which for me is more than twice the cost of
the items additionally ). This may be all to no avail if this doesn't work as

The cheapest retail compares to bulk 50 lb prices
MnSO4 H2O - 5 lb - 5.99
also of interest there _
(NH4)2SO4 - 20 lb - 8.99

Twice as expensive compared to the source above but the nitrate is
as cheap as I have seen , and there is also urea available but not so cheap.
MnSO4 H2O - 4 lb - 9.90
NaNO3 - 4 lb - 3.62
H2NCONH2 - 5 lb - 12.38
(NH4)2SO4 - 20 lb - 8.55

A local green nursery should be the source lowest in overall cost.

[Edited on 23-10-2009 by franklyn]

kilowatt - 23-10-2009 at 11:04

Oh yeah, I read that wrong, being as MgSO4 is a lot more common than MnSO4 my mind just sort of extrapolated the n into a g. Sorry about the confusion.

In any case, I do find $5 for 5lbs of MnSO4 to be expensive, at least with regards to sulfate molarity. However MnO2 is cheap from ceramics suppliers, so perhaps synthesizing it from this would be a good alternative, or at least good for recycling the MnO2 left after the reaction. Bisulfite salts or another SO2 source may be an economical recycling method since I believe MnO2 would react with them to make the (II)sulfate.

While we're on the subject of a synthesis that already involves heating things to high temperatures, I have obtained good yields of SO2 by reducing common sulfates with carbon at ~1000°C, so this would be a way to recycle the Na2SO4 that is left after the reaction. I believe leading the hot SO2 through MnO2 powder in another vessel would suffice, and would be the ideal method if sodium nitrate is your nitrate source. If the MnO2 could itself catalyze the reduction, it may be possible to lower the temperature enough to allow the overall reaction Na2SO4 + MnO2 + C --> MnSO4 + CO + Na2O to occur at once in a solid cake, eliminating the need even for a retort or any sort of vessel. If that were the case, that final solid product could be added to NH4NO3 to produce the reaction mixture of interest. Looks like this calls for some experimentation.

[Edited on 23-10-2009 by kilowatt]

franklyn - 23-10-2009 at 12:52

Before anything else I have corrected the amount of product I gave above :
13 mols of N2O4 ( 92 wt.) is 1196 grams not 1612 - I did not account for
the oxygen remaining with MNO2 ( 13 mols of O2 ( 32 wt.) is 416 )

@ killwatt
You really want to sqeeze every drop of economic value from this !
Consider that nearly ~ 1200 grams at ~ $ 10 + , is economical even at rocketry
scale , ~ $6 in 1990 -

How much is it now ? You can always price it , I doubt you could take delivery though.

NH4NO3 + MnSO4 -> ? , I don't know , decomposition of NH4NO3 occurs at 210 ºC
and that of (NH4)2SO4 at under 280 ºC remember dry distillation is the point.
How to keep MN(NO3)2 dry. That would invoke the use of non aqueous solvents.


kclo4 - 23-10-2009 at 14:47

I like this idea, but also consider mixing and heating Na2S2O7 and NaNO3.
That should form Na2SO4 and N2O4.

Na2S2O7 might be easier to obtain then MnSO4 to some members, so I figured I'd mention it.
I know it is easier for me to get.

497 - 23-10-2009 at 16:22

That should form Na2SO4 and N2O4.

The reaction doesn't balance unless you account for the additional O2 formed.. Which causes serious problems when attempting to condense the N2O4... But if you are just bubbling it into water to make HNO3, or something else where O2 isn't a problem, then yes I agree it may be as suitable, or even preferable to the proposed manganese route.. But personally my main interest is in condensing near 100% pure N2O4..

Na2SO4 + MnO2 + C --> MnSO4 + CO + Na2O to occur at once in a solid cake, eliminating the need even for a retort or any sort of vessel. If that were the case, that final solid product could be added to NH4NO3 to produce the reaction mixture of interest.

Hmmm I like it. The overall reaction being 2NH4NO3 + C --> 2NH3 + N2O4 + CO + H2O. And the simplicity is wonderful. Just pyrolyze off the N2O4, add carbon to reform MnSO4 and Na2O, add NH4NO3 to regenerate the NaNO3, and repeat...

Except.. It seems to me that the MnO2 + Na2SO4 + C reaction may not work as you stated, but will just result in MnO and/or Mn2O3 mixed with Na2SO4 (or Na2O if more C was added). I see no reason for the SO4-- to transfer to the Mn++ rather than stay with the Na+.. But that problem could be avoided by simply running the standard reduction of Na2SO4 by C and absorbing the SO2 into MnO2. Only a little added complexity..

It would be nice if the SO2 would effectively absorb into dry MnO2 powder, so that water is not involved.. But I suspect you might have a hard time getting a complete conversion of the MnO2 that way. Literature on the subject states that often only 10-50% of the MnO2 is reacted.. Although they're usually talking about SO2 concentrations on the order of <1%.. So maybe at high concentrations it would react more completely. If the dry reaction failed, a suspension of MnO2 is known to easily and quantitatively react to form a MnSO4 solution.. It would just have the added energy intensive step of drying the MnSO4.

Although it is a very interesting possibility, I see no reason to go to the effort of dealing the 1000*C reactions and bulk CO, and NH3, when you could simply use K/NaNO3 as your NO3- source, throw away your waste sulfate, and regenerate the the MnO2 by burning sulfur (or for that matter, buying MnSO4...) Yes, of course recycling more of the chemicals will save money and allow cheaper feedstocks, but would it really be worth doing on anything less than a truly massive scale? After all, we're talking about differences of a few dollars per Kg of N2O4 here...

[Edited on 24-10-2009 by 497]

kilowatt - 23-10-2009 at 18:34


NH4NO3 + MnSO4 -> ? , I don't know...

I was accounting for the Na2O left after the carbon reduction of Na2SO4 reacting with the NH4NO3 to produce Na2NO3 and liberating ammonia, not a direct reaction of MnSO4 and NH4NO3 at high temperature. Unless there would be a detrimental reaction between those things at room temperature? I don't think there would be though.

Anyways it is true that recycling in this manner would only pay off at rather large scale, and if you had use for either excess Na2O or NH3 (depending on your nitrate source) and could use or dispose of the CO. What scale that might be I'm not sure, but I imagine something like tens of kg of N2O4. It could be an interesting process anyway, and a possible way to produce more than one useful chemical at once from other chemicals. Admittedly more work than I'm going go for in the near future.

Regarding the incomplete reaction of SO2 and dry MnO2, I had considered this, and it would not be much of a problem if all the MnO2 was not reacted; it could simply be separated by dissolving in water and filtering, but of course that would require additional energy to dry it out as mentioned. Unreacted SO2 would still be a problem. Perhaps it could be recycled or run through a large excess of MnO2. Maybe a fluidized bed reactor would be the way to go, but that's fairly elaborate. An aqueous suspension or suspension in some more volatile solvent sounds most plausible to me.

[Edited on 24-10-2009 by kilowatt]

kclo4 - 23-10-2009 at 18:53

From my understanding, it is best to soak the MnO2 in a sulfuric acid solution to get it to react with the SO2 to form MnSO4.

franklyn - 24-10-2009 at 18:53

I did emphasize cheap in the thread topic heading for this reason ,
in bulk it just keeps getting cheaper , hardly worth ' recycling '
when the by product is not even hazardous waste. See _

Manganese Sulfate , 50 lb - $ 50.90
Chilean Nitrate , 50 lb - $ 51.70 , * this can be had cheaper

This much will yield just short of 12 kilograms N2O4 , at $ 8.55 per Kg or $ 47 a gallon
comparable to retail cost of nitromethane or methylene chloride and common solvents.
The cost per unit weight is almost the same as for the 5 lb bags I gave above.

Another interesting source for Manganese Sulfate , 50 lb - $ 41.35
Also , Nitric Acid 165 lb. drum 67.2% - $ 22.99 , Magnesium Nitrate , 50 lb - $ 39.99
Rosco Bodine posted here how Mg(NO3)2 can be used to distill HNO3 to highest conc.


kilowatt - 24-10-2009 at 20:57

That last supplier looks almost too good to be true, super low prices, wide variety of "real" chemicals. Some of those things (bulk quantities of nitric acid, sulfuric acid, sulfur, nitrates) don't seem like things that an individual could normally buy, at least without expecting a BATFE or DEA raid. In any case I hate to think what shipping would be on some of those items, especially with the hazmat on the corrosive liquids. Will these people really sell to an individual no questions asked? Has anyone tried (maybe with some fairly innocuous stuff first)? Their mono-ammonium phosphate would be an awesome buy for phosphate bonded investment casting mixtures.

[Edited on 25-10-2009 by kilowatt]

franklyn - 27-10-2009 at 10:58

Not to bump my own thread again , it seems some are dismissive of
the risk - emphasis must be on safe practice when producing NO2 ! !

Health concerns



franklyn - 17-1-2010 at 20:58

I found this precedence in prior research

Farnell, Westerdahl and Taylor, 1972 , Third International Pyrotechnics Seminar
The Influence of Transition Metal Compounds on the Aluminum Sodium Nitrate Reaction
Manganese Sulfate MnSO4•H2O , added to Aluminum Sodium Nitrate flare compositions ,
alters " the decomposition of Sodium Nitrate to form oxides of nitrogen rather than its
normal decomposition products of nitrogen and oxygen.

perhaps someone can obtain and post acopy of the paper.

* Note that some MSDS cite dehydration temperature for anhydrous MNSO4
at 400 - 500 ºC. If this is so then heating will need to be done stove top as
the oven will not provide that high temperature.

[Edited on 18-1-2010 by franklyn]

Paddywhacker - 17-1-2010 at 22:04

Iron II as in ferrous sulphate would also be expected to absorb an oxygen, to become ferric oxide.

But heating FeSO4.7H2O will release H2SO4 and that will react with the NaNO3 to produce HNO3 fumes. So, unless you can get a fusable nonvolatile anhydrous iron II salt, then iron is out.

franklyn - 1-12-2010 at 05:44

Found that reference :

Third International Pyrotechnics Seminar ( 1972 ) pg 271
The Influence of Transition Metal Compounds on the Aluminum Sodium Nitrate Reaction
Farnell , Westerdahl , Taylor , The reaction of interest is the first on pg 285
redirects to :

An attached excerpt from the Handbook of Solvents
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cited references are here

Molten salt solutions.gif - 87kB