teodor
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Anhydrous nitrates: N2O4 way
It was a nice early autumn evening.
Already several days (may be weeks or months, the time in a lab going very fast) I did experiments with sodium hexametaphosphate mixtures (this is not
a single compound but a mixture of condensed phosphates in a form of rings with different number of elements, 6 and 5 are most common).
This particular series was about studying the reactions of the general type
4 NaNO3 + 4 NaPO3 -> 2 Na4P2O7 + 4 NO2 + O2
under different conditions. It is molten-state reaction (at least one salt must be molten or at least soften), so for heating I prepared the Rose’s
metal by melting together 2 parts of Bi, 1 part Sn and 1 part Pb:
I use a soldering pot as a heater which allows set bath temperature in the range 100 - 450C with the precision of 1 degree Celsium.
Because in this series I deal with NO2 I need an effective scrubber. As my previous experiments revealed (https://www.sciencemadness.org/whisper/viewthread.php?tid=15...) concentrated H2SO4 is a very good scrubber both for NO and NO2 but when I tried
to use it in the bottom flask it was not effective because as I realized NO2 should be bubbled through it. So I’ve attached a Dimroth condenser to
condense NO2 to N2O4 first and I thought that when drops of N2O4 will fall from the condenser into H2SO4 it would be perfectly adsorbed. The problem
appears, it is not possible to condense NO2 with this type of condenser. The spiral one must be used:
To the 2nd neck of the bottom flask I’ve attached Dimroth condenser (just to lift the line) and than I put 2 scrubbers: P2O5 (it forms adducts with
NO on heating) and H2SO4.
Several experiments was very successful by means of condensing N2O4 and mixing it with H2SO4. For the coolant I used ethanol cooled with the immersion
cooler (it is behind the desiccator on the photo) cooled to -5C. I was afraid about the spiral clogging but eventually N2O4 is very mobile and could
be slightly overcooled. But -5 is enough (the b.p is 21C, the m.p. is -11).
When the first drops of N2O4 went through the condenser very interesting changing of color was observed. The liquid N2O4 is yellow because of
dissolved NO2. Some organic residue was in the spiral, so going down the color was changing to intense green (NO2 / N2O3 in N2O4), then light blue and
at the end colorless. This colorless was probably a pure N2O4 after reducing and reoxidyzing of NO2. It looks like N2O4 is more inert as to organic,
so it doesn’t change color when all dissolved NO2 reacts (I have some video of this process on youtube).
But the goal was achieved. Initially 16 ml of H2SO4 around 90% concentration was charged in the bottom flask but after collecting N2O4 from different
experiments (and apparently some HNO3) the volume becomes 40 ml. It becomes straw-colored fuming liquid with a strong odor of HNO3.
Then I tried to collect N2O4 without H2SO4 adsorbent but the vapor density around 20.5 C (the lab temperature) it too high, so part of it stays in the
liquid phase but part as NO2 vapor.
So, I put 20 ml of ethyl acetate. Ethyl acetate is a good solvent for N2O4 (C. C. Addison and N. Logan "Anhydrous Metal Nitrates" in "Advances in
Inorganic Chenistry and Radiochemistry", vol. 6)
it becomes donor for electon-deficient molecule of N2O4 and the equilibrium is set:
[(Don)n * NO]+ + NO3- <-> n(Don) + N2O4 <-> (Don)n * N2O4.
The ionization to NO+ and NO3- is very important here. Actually, it is the action of the solvent. (N2O4 is good in making different types of
dissociation, really 3 types: NO2 + NO2, NO2- + NO2+ and NO+ + NO3-, so this particular solvent activates the 3rd type).
I checked stability of ethyl acetate in the steam of NO2 (dropping it into the heated flask) up to 360C metal bath temperature. There is some
nitrate/nitrile smell after opening the apparatus, but no violent reaction happens. As to NO2 adsorbing from a system it is almost as good as water,
it can scavenger NO2 in a vapour form.
Trying to utilize the mixture I put some cobalt metal in a form of plates. It reacts quite actively forming deep red solution.
On the bottom there are some crystals and the unreacted metal:
I think it is possible to crystallize 2 types of compounds: nitrosonium nitrocobaltate(II) and anhydrous cobalt nitrate. (See e.g. Tikhomirov et al
"Anhydrous Nitrates and Nitrosonium Nitrometallates of Manganese and Cobalt", Z. anorg. allg. Chem. 628 (1): 269-273). By the color I think this one
is a nitrate. Also because there is some unsolved cobalt. For nitrosonium salt the excess of N2O4 is required.
I put my solution in a freezer and made a photo of deposited crystalls:
This is my first experiment of successful NO2 condensation, dissolving it in a solvent and getting an anhydrous nitrate of transition metal.
Any comments and suggestions are appretiated.
[Edited on 11-9-2025 by teodor]
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woelen
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This really is interesting. I have read about anhydrous copper nitrate from ethyl acetate, but never attempted to try this myself. Good to see that
you get this result.
What surprises me is the bright pink color of this cobalt compound. I would expect a blue color. I have done quite some experiments with cobalt in the
past, and as soon as water is involved, it becomes pink or rose colored. Anhydrous complexes nearly always tend to be blue. You can even get anhydrous
complexes of cobalt in aqueous solution, e.g. if you drop a few crystals of CoCl2.6H2O in conc. HCl, then you get a deep blue solution. IIRC, a
complex CoCl4(2-) is formed. On dilution, the liquid turns pink.
Do you have a means for checking water content of your sample? Could it be a mixed complex, also containing some water?
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teodor
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I didn't check the water contents yet becaus I have no clear idea how to do it, so I probably will delay it keeping the liquid for the future
experiments. What I did, I tried to dissolve hexahydrate in ethyl acetate. You can do it by dropping crystalls and melting hydrate. 2 layers are
formed and then diffusion happens, in the organic layers the salt is almost immediatelly deposited in the form of very pale good formed crystalls,
after few hours the water layer disappeares. The crystalls are very similar indeed. So I unable to decide. I can say that for this experiment I took
dry salts and no water vapours were visible in a glass pipe which connects the flask and the condenser. But if it was in a form of HNO3 it could not
be visible.
But I did a check of publications Co(NO3)2 is pink-purple, (NO)2[Co(NO3)4] is red-violet.
https://doi.org/10.1002%2F1521-3749%28200201%29628%3A1%3C269...
https://doi.org/10.1002/cber.19640970536
The wikipedia article gives "pale red powder".
What I can relatively easily to do is to prepare another transition metal nitrate by the similar method to check whether my salt mixture gives HNO3.
This question is important for my experiments, it means that there are some acidic content (H) in the salt. So I need to find a transition metal which
will react with N2O4 and will give completely different colors for the hydrate and anhydrous forms. Any suggestions?
P.S. The ehyl acetate could be not perfectly dry also.
P.P.S. I know it is dependent on my monitor, but I would say I've got #7a1a30
I have a mixture of 17000K and 6500K lamps in my laboratory, so the color balance is more cold than sunlight. This is the photo with incandenscent
light:
And at the daylight with the color index:
[Edited on 11-9-2025 by teodor]
[Edited on 11-9-2025 by teodor]
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DraconicAcid
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Quote: Originally posted by woelen  | This really is interesting. I have read about anhydrous copper nitrate from ethyl acetate, but never attempted to try this myself. Good to see that
you get this result.
What surprises me is the bright pink color of this cobalt compound. I would expect a blue color. I have done quite some experiments with cobalt in the
past, and as soon as water is involved, it becomes pink or rose colored. Anhydrous complexes nearly always tend to be blue. |
The pink isn't due to the water, but due to an octahedral geometry instead of tetrahedral. I've also read about the anhydrous copper(II) nitrate from
ethyl acetate, and I seem to recall that the nitrate is at least partially bidentate. That could make the cobalt octahedral.
Please remember: "Filtrate" is not a verb.
Write up your lab reports the way your instructor wants them, not the way your ex-instructor wants them.
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Fery
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Great experiment and photos with beautiful colors.
I once saw liquid N2O4, the chemist told me that he made it by himself by decomposition of Pb(NO3)2 and the hardest part was efficient cooling +
condenser. He used few condensers in a cascade. First stage air condenser. Final stage Graham condenser, salt+ice for cooling. Circulated liquid in
the last stage condenser had about 0 C or only slightly below (salty water). Collecting flask submerged in ice+salt bath. Slow heating so slow ratio
of the product formation.
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teodor
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| Quote: Originally posted by Fery | Great experiment and photos with beautiful colors.
I once saw liquid N2O4, the chemist told me that he made it by himself by decomposition of Pb(NO3)2 and the hardest part was efficient cooling +
condenser. He used few condensers in a cascade. First stage air condenser. Final stage Graham condenser, salt+ice for cooling. Circulated liquid in
the last stage condenser had about 0 C or only slightly below (salty water). Collecting flask submerged in ice+salt bath. Slow heating so slow ratio
of the product formation. |
Yes, the success of liquidification is partially because the speed of decomposition by means of hexamethaphosphate is very easily to control. This is
surface-to surface reaction between two salts, not the single salt decomposition. I adjust the temperature until I have a constant flow of N2O4 in the
condenser. There are drops in every turn of a spiral going down. The actual condensation after some thermal equilibrium is at the top of the
condenser. I think Allihn type can also work for ascending vapours. Those constrictions are essential, the condensation happens on constrictions.
By the way, for drying salts I use the dessicator you gave me once. I replaced the lid to be able to use it with vacuum.
N2O4 could be dissociated NO2+ and NO2-, so I think should be a way to use it as a nitration mixture. Not experimenting with this yet. But I ordered
few more organic solvents and will try modify N2O4 behaviour with different solvent.
Thank you very much, Fery.
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teodor
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I've just made a very rough test putting 2 drops of water into the reaction flask.
And this is chemically pure hexahydrate for comparison:
 
Still, the question is it (NO)2[Co(NO3)4] or Co(NO3)2 remains.
P.S. This pink residue in the flask corresponds to #ed5a84 on my monitor. Probably there is a difference between dry and wet matherial.
[Edited on 11-9-2025 by teodor]
[Edited on 11-9-2025 by teodor]
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teodor
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Surely, that cobalt chemistry is colourful. But. N oxides hemistry is not less colorful!
I've made a picture of reacting N2O4/NO2 with some compound (I will check what it was, need to check how I used this condenser before) and put my
interpretation of the color changes.
So, I am quite sure there is a way to get liquid colourless N2O4 with sime kind of reducer-stabilizer.
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chloric1
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Amazing work! Nothing like enjoying autumn nights after an especially hot and muggy summer! I love how you use the hexametaphosphate as an acid to
liberate NOx! BTW the color of cobalt nitrate hexahydrate makes sense as it looks very much like the hydrated cobalt sulfate I recently made.
That’s why I’m partial to cobalt chloride because of the sky blue anhydrous salt and the range of purples and pinks with hydrated chlorides.
Fellow molecular manipulator
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teodor
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| Quote: Originally posted by chloric1 | | Amazing work! Nothing like enjoying autumn nights after an especially hot and muggy summer! I love how you use the hexametaphosphate as an acid to
liberate NOx! BTW the color of cobalt nitrate hexahydrate makes sense as it looks very much like the hydrated cobalt sulfate I recently made.
That’s why I’m partial to cobalt chloride because of the sky blue anhydrous salt and the range of purples and pinks with hydrated chlorides.
|
Thank you chloric1! We didn't have particulaly muggy summer, but I like autumn anyway.
Not exactly like an acid. It gets two cations and one oxygen. When it reacts e.g. with NaCl (NaBr, NaI) it consumes O2 from air. And there is no
hydrogen involved. And the reaction product is not Cl-, but Cl2 (Br2, I2). So, there are important differences but many reactions could be made by
this scheme as we usually do with acids but with those differences just mentioned. The limitations is a temperature range and the requirement of some
reaction media - molten salt or a liquid which is not evaporated in the active region of the temperature range. Also, hydrogen can change something.
But you can imagine a lot of reactions of the general scheme as Me2O/H2O elimination. The question is adjusting conditions to make them happen.
My study now is about the "active region" (the range of temperatire) when some particular composition works. If we know this we can select the media
as a molten salt, acid, watever. But the first question is the active region.
The cobalt is the nice metal. His chemical character is more "pronounced" than that of chromium and more colorful than copper for transition metal
chemistry. I made sulfate from Co + H2SO4, the well formed big crystalls was made. Also I made different attempts to dehydrate chloride (have some
nice pictures: http://www.sciencemadness.org/talk/viewthread.php?tid=78561&...). Here also: https://www.sciencemadness.org/talk/viewthread.php?tid=78561...
But the most interesting for me was preparing some complexes which could be used to prepare another complexes, so for each complexes there is some
path how to jump from one compound to another. I did only beginning of this story. Some compounds have no reported explosiveness, so I did some
experiments with this also: https://www.sciencemadness.org/talk/viewthread.php?tid=15861...
I think there could be also some other my publications about cobalt. May be about some Co(III) chellate and the 3rd complex from the first story. I
will be glad to revisit that but that is quite well studied field (to start you can use publications in the Inorganic Syntheses series) unlike those
metaphosphates were there are many things still not explored or unpublished (or published in some journals nobody can find).
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chloric1
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Well one aspect I like about your approach is I could make cupric nitrate anhydrous and VERY carefully add the required water to form the trihydrate
instead of trying to wait 2+ months for a concentrated solution to crystallize! I like to make my own patina solutions to color metals. On brass and
copper cupric nitrate can form greens or blues based on what other copper salts or alkaline salts are present. Cupric nitrate even attacks zinc to
form a greyish black color. I agree with you on the colorfulness of cobalt chemistry. Chromium is nice but a little hard to work with unless you’re
making chromates and dichromates. The sulfato chromium complexes can be a bitch and at least a part of them would not even precipitate barium sulfate
from solutions of barium salts!
I think the Netherlands would be an awesome place to visit, especially in summer. Where I live we max out 95-100 Fahrenheit(35-38 Celsius) often with
60% humidity! The stifling heat is a tall order considering we dip to -20 ℃ in January!
Fellow molecular manipulator
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teodor
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| Quote: Originally posted by chloric1 | Well one aspect I like about your approach is I could make cupric nitrate anhydrous and VERY carefully add the required water to form the trihydrate
instead of trying to wait 2+ months for a concentrated solution to crystallize!
|
Try this: take a small part of the solution and seed it with a rhombohedral crystal of other salt. Of course some other metal nitrate hydrate should
work the best way. But the key point that isomorhphic crystals cause coprecipitation. The salt doesn’t care the seed crystals composition, only
geometry when crystallize. So, you can try to generate seed crystals this way. This is generally how the chicken-egg problem for some crystals was
ever solved.
I would give priority to nitrates and halides as seeds. I could be wrong, but it seams some halogen have a similar ion size to nitrate.
For this particular salt it is good to add a bit of HNO3 to prevent turning to basic nitrate on precipitation (Perin, Armarego and Perin
“Purification of laboratory chemicals”)
[Edited on 13-9-2025 by teodor]
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chloric1
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| Quote: Originally posted by teodor | | Quote: Originally posted by chloric1 | Well one aspect I like about your approach is I could make cupric nitrate anhydrous and VERY carefully add the required water to form the trihydrate
instead of trying to wait 2+ months for a concentrated solution to crystallize!
|
Try this: take a small part of the solution and seed it with a rhombohedral crystal of other salt. Of course some other metal nitrate hydrate should
work the best way. But the key point that isomorhphic crystals cause coprecipitation. The salt doesn’t care the seed crystals composition, only
geometry when crystallize. So, you can try to generate seed crystals this way. This is generally how the chicken-egg problem for some crystals was
ever solved.
I would give priority to nitrates and halides as seeds. I could be wrong, but it seams some halogen have a similar ion size to nitrate.
For this particular salt it is good to add a bit of HNO3 to prevent turning to basic nitrate on precipitation (Perin, Armarego and Perin
“Purification of laboratory chemicals”)
[Edited on 13-9-2025 by teodor] |
Yeah a little free nitric acid goes a long way in keeping heavy metal nitrates stable! This what is necessary for lead and ferric nitrate as well.
Of course ferric ions being smaller and having high charge density make it considerably more covalent so more nitric acid is needed.
Fellow molecular manipulator
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