Rhodanide
Hazard to Others
Posts: 348
Registered: 23-7-2015
Location: Ｔｈｅ ８０ｓ
Member Is Offline
Mood: That retro aesthetic
|
|
Making the bigger brother of Chloropicrin - is the rxn. going as planned?
So!
Last week I took a shot at making liquid N2O4. Long story short, my condenser wasn't cold enough and it didn't work. Not being one to completely give
up, I recalled the preparation of 1,1,2,2-tetrachloro-1,2-dinitroethane. This required "Nitrogen Peroxide" (Essentially NO2) to be sealed in a
container with Tetrachloroethylene*, which I happened to have on hand, and the mixture heated for an hour. (*Note: This was obtained from a certain
aerosol product for cleaning, and contained only a tiny, tiny amount of rust inhibitor, which made it pure enough for my purposes)
I ended up just bubbling NO2 through the TCE for a good 10 minutes. After this time, I was left with a strikingly grass-green solution in a red-brown
atmosphere of NO2. It looked quite pretty. I never ended up getting to heat it, due to other complications concerning time, so I left it stoppered in
a reasonably cool place. Thing is, the literature never stated anything about a green solution. The lit. stated that the mixture was allowed to
evaporate off the NO2 and residual TCE, leaving a white residue which is presumably the product. Is NO2 soluble in TCE, and if so, why is it green? Is
there a side reaction I'm not aware of? Do any of you reckon I need to heat it to get the desired product?
[Edited on 1-2-2017 by Tetra]
|
|
|
woelen
Super Administrator
Posts: 8027
Registered: 20-8-2005
Location: Netherlands
Member Is Offline
Mood: interested
|
|
Your NO2 must have been (somewhat) impure. It also contained NO. When you dissolve NO2 in a solvent, then the color is brown/red. At low temperature
(around 0 C) the color of the dissolved NO2 is pale brown/yellow, most of it then forms N2O4, which is colorless. In the presence of NO, however, you
get N2O3 in your solution. N2O3 is a liquid with a deep blue color. Mixed with the red/brown/yellow colors of NO2+N2O4 this gives a green color.
How did you make your NO2? Making it from HNO3 and some reductor always leads to impure NO2. Even if you use warm concentrated acid, then still there
is some NO in the gas, which on cooling down leads to formation of a dark blue or dark green liquid. You can avoid this problem if you have a second
gas generator, which makes oxygen and if you lead a stream of oxygen into the stream of NO2+NO and then allow the mix to flow through a thin tube
before it goes into the C2Cl4. The NO is quickly oxidized by the O2 and if you have excess O2, then you get a mix of NO2 and O2. In your application
the excess O2 is no problem, it simply bubbles through the solution.
|
|
|
Rhodanide
Hazard to Others
Posts: 348
Registered: 23-7-2015
Location: Ｔｈｅ ８０ｓ
Member Is Offline
Mood: That retro aesthetic
|
|
Quote: Originally posted by woelen  | Your NO2 must have been (somewhat) impure. It also contained NO. When you dissolve NO2 in a solvent, then the color is brown/red. At low temperature
(around 0 C) the color of the dissolved NO2 is pale brown/yellow, most of it then forms N2O4, which is colorless. In the presence of NO, however, you
get N2O3 in your solution. N2O3 is a liquid with a deep blue color. Mixed with the red/brown/yellow colors of NO2+N2O4 this gives a green color.
How did you make your NO2? Making it from HNO3 and some reductor always leads to impure NO2. Even if you use warm concentrated acid, then still there
is some NO in the gas, which on cooling down leads to formation of a dark blue or dark green liquid. You can avoid this problem if you have a second
gas generator, which makes oxygen and if you lead a stream of oxygen into the stream of NO2+NO and then allow the mix to flow through a thin tube
before it goes into the C2Cl4. The NO is quickly oxidized by the O2 and if you have excess O2, then you get a mix of NO2 and O2. In your application
the excess O2 is no problem, it simply bubbles through the solution. |
Made it by reacting NaNO2 and dilute H2SO4.
|
|
|
Praxichys
International Hazard
Posts: 1063
Registered: 31-7-2013
Location: Detroit, Michigan, USA
Member Is Offline
Mood: Coprecipitated
|
|
Yes, you absolutely need heat and pressure for this reaction to work. Your heating method is ineffective; the solubility of a gas in a liquid
decreases with temperature. All you did by heating it for an hour is get rid of some of the NO2. I agree with woelen on the cause of the blue
coloration. NO contamination could have also been created in-situ with the reductive properties of the rust inhibitor, though it is most likely
impurities from the way in which it was prepared.
That said, the conditions given in this preparation involve heating a closed container with two volatile liquids to 100°C which puts this reaction
under modest pressure. NO2/N2O4 should be under about 280 PSI at 100°C, FYI. I doubt this reaction occurs to any extent at atmospheric pressure, so
you will need a suitable pressure vessel.
The real danger here is that NO2/N2O5 above about 150°C starts to decompose into nitrogen and oxygen, which (although an endothermic process) will
certainly cause the pressure to spike, potentially causing an explosion. Also, the evolved oxygen may be reactive enough at that temperature and
pressure to produce phosgene from the TCE. Careful temperature control is required; I highly recommend a boiling water bath if you attempt this at
all.
https://cameochemicals.noaa.gov/chris/NOX.pdf
|
|
|
Rhodanide
Hazard to Others
Posts: 348
Registered: 23-7-2015
Location: Ｔｈｅ ８０ｓ
Member Is Offline
Mood: That retro aesthetic
|
|
| Quote: Originally posted by Praxichys | Yes, you absolutely need heat and pressure for this reaction to work. Your heating method is ineffective; the solubility of a gas in a liquid
decreases with temperature. All you did by heating it for an hour is get rid of some of the NO2. I agree with woelen on the cause of the blue
coloration. NO contamination could have also been created in-situ with the reductive properties of the rust inhibitor, though it is most likely
impurities from the way in which it was prepared.
That said, the conditions given in this preparation involve heating a closed container with two volatile liquids to 100°C which puts this reaction
under modest pressure. NO2/N2O4 should be under about 280 PSI at 100°C, FYI. I doubt this reaction occurs to any extent at atmospheric pressure, so
you will need a suitable pressure vessel.
The real danger here is that NO2/N2O5 above about 150°C starts to decompose into nitrogen and oxygen, which (although an endothermic process) will
certainly cause the pressure to spike, potentially causing an explosion. Also, the evolved oxygen may be reactive enough at that temperature and
pressure to produce phosgene from the TCE. Careful temperature control is required; I highly recommend a boiling water bath if you attempt this at
all.
https://cameochemicals.noaa.gov/chris/NOX.pdf
|
I haven't officially begun the heating yet. 
That's just what the lit said to do.
|
|
|
woelen
Super Administrator
Posts: 8027
Registered: 20-8-2005
Location: Netherlands
Member Is Offline
Mood: interested
|
|
That produces a mix of NO and NO2:
2 NaNO2 + 2H(+) --> 2 HNO2 + 2 Na(+)
followed by:
2 HNO2 <--> N2O3 + H2O
followed by:
N2O3 <--> NO + NO2
If you add dilute H2SO4 to NaNO2 then you can see the blue color of N2O3 in the solution. The color never becomes very dark, due to quick
decomposition in water, but still it is visible well, the liquid becomes pale blue, like a fairly diulte solution of CuSO4.
|
|
|
Rhodanide
Hazard to Others
Posts: 348
Registered: 23-7-2015
Location: Ｔｈｅ ８０ｓ
Member Is Offline
Mood: That retro aesthetic
|
|
| Quote: Originally posted by woelen | That produces a mix of NO and NO2:
2 NaNO2 + 2H(+) --> 2 HNO2 + 2 Na(+)
followed by:
2 HNO2 <--> N2O3 + H2O
followed by:
N2O3 <--> NO + NO2
If you add dilute H2SO4 to NaNO2 then you can see the blue color of N2O3 in the solution. The color never becomes very dark, due to quick
decomposition in water, but still it is visible well, the liquid becomes pale blue, like a fairly diulte solution of CuSO4. |
That was observed indeed, but I thought the blue color itself was due to the HNO2 as a chemical. Is it the N2O3 or is the HNO2 really blue?
I ended up heating the blue mixture until no more NOx came off. It was clear at the end.
I'll just give the link to the synth. http://www.prepchem.com/synthesis-of-1-1-2-2-tetrachloro-12-...
[Edited on 2-2-2017 by Tetra]
|
|
|
woelen
Super Administrator
Posts: 8027
Registered: 20-8-2005
Location: Netherlands
Member Is Offline
Mood: interested
|
|
N2O3 is blue, HNO2 is colorless.
A nice demonstration of that is adding NaNO2 to conc. H2SO4. In that case you don't see any blue color. You get a colorless solution, which contains a
mix of HNO2 and NO(+) ions:
NaNO2 + H2SO4 --> HNO2 + NaHSO4
followed by the equilibrium reaction
HNO2 + H2SO4 <--> NO(+) + H2O + HSO4(-)
|
|
|
PHILOU Zrealone
International Hazard
Posts: 2893
Registered: 20-5-2002
Location: Brussel
Member Is Offline
Mood: Bis-diazo-dinitro-hydroquinonic
|
|
Better start from 1,2-dinitroethan (DNE) and halogenate in basic media to get tetrahalogeno-dinitroethan.
DNE can be obtained via:
-addition of N2O4/NO2 to ethylen (forms also side products nitrite ester, nitrate ester, oxydation products, nitroso, nitro or mix of those)
-addition (substitution) of NaNO2 to 2,3-dihalogeno-butan-1,4-dioic acid (dihalogeno-succinic acid)...mimicking the formation of nitromethane from
chloroacetic acid with NaNO2.
The introduction of a nitro group favourize the elimination of the carboxylic moeity from the carbon holding the NO2 group.
The NO2 may enter as nitrite ester and lead to oxydation products or nitrosation products of the nitro compound (nitronic acids)...urea may reduce the
nitrite ester troubles and side products and increase nitrite salt solubility.
-addition (substitution) of AgNO2 or Na/K/LiNO2 to 1,2-dihalogeno-ethane into a suitable solvent (DMF or DMSO) with urea to reduce the nitrite ester
troubles and side products and increase nitrite salt solubility.
Possible formation of nitro-nitrite ester or dinitrite ester aside from DNE.
[Edited on 4-2-2017 by PHILOU Zrealone]
PH Z (PHILOU Zrealone)
"Physic is all what never works; Chemistry is all what stinks and explodes!"-"Life that deadly disease, sexually transmitted."(W.Allen)
|
|
|
Boffis
International Hazard
Posts: 1879
Registered: 1-5-2011
Member Is Offline
Mood: No Mood
|
|
Hi Philou, do you have a reference for the reaction of 2,3 dihalo-butan-1,4-dioic acid with sodium nitrite? I made some of the dibromo acid from
bromine and fumaric acid (1) as a source for butyn-1,4-dioic acid (2) which was intended to be a stepping stone towards dicyanoacetylene via the
diester then diamide but never got round to finishing it off because I discovered a paper that described precisely what I was intending and it failed
(Dicyanoacetylene + NaN3 + acid -/-> 1,2 bis(5-tetrazolyl)acetylene). 1,2 dinitroethane sound like an interesting compound.
1) Dibromosuccinic acid from fumaric acid; Org. Synth.; CV2 p177
2) Acetylenedicarboxylic acid; Org. Synth.; CV2 p10
|
|
|
PHILOU Zrealone
International Hazard
Posts: 2893
Registered: 20-5-2002
Location: Brussel
Member Is Offline
Mood: Bis-diazo-dinitro-hydroquinonic
|
|
| Quote: Originally posted by Boffis | Hi Philou, do you have a reference for the reaction of 2,3 dihalo-butan-1,4-dioic acid with sodium nitrite? I made some of the dibromo acid from
bromine and fumaric acid (1) as a source for butyn-1,4-dioic acid (2) which was intended to be a stepping stone towards dicyanoacetylene via the
diester then diamide but never got round to finishing it off because I discovered a paper that described precisely what I was intending and it failed
(Dicyanoacetylene + NaN3 + acid -/-> 1,2 bis(5-tetrazolyl)acetylene). 1,2 dinitroethane sound like an interesting compound.
1) Dibromosuccinic acid from fumaric acid; Org. Synth.; CV2 p177
2) Acetylenedicarboxylic acid; Org. Synth.; CV2 p10 |
I don't have a reference for it.
Even if a bit more sterically stressed than haloacetic acid because the halide is secondary and not primary (nitrite substitution on secondary carbon
is harder but works);
it should work because the halide atom is very strongly activated by the alfa carboxyl and halide and by the beta carboxyl.
Also once the first halide is replaced by a NO2; then the second one is even more activated by the new alfa nitro group.
Such dibromosuccinic acid is also very succeptible to hydrolysis by NaOH to deliver tartric acid.
Here is an example of diamination of dibromosuccinic acid with large amine (allylamine) and it works:
CLXXXII.—The action of aliphatic amines on s-dibromosuccinic acid. Part II. Allylamine
Edward Percy Frankland and Henry Edgar Smith
J. Chem. Soc., Trans., 1912,101, 1724-1729
DOI: 10.1039/CT9120101724
2 CH2=CH-CH2-NH2 + HO2C-CHBr-CHBr-CO2H --> CH2=CH-CH2-NH-CH(CO2H)-CH(CO2H)-NH-CH2-CH=CH2. 2 HBr
[Edited on 5-2-2017 by PHILOU Zrealone]
PH Z (PHILOU Zrealone)
"Physic is all what never works; Chemistry is all what stinks and explodes!"-"Life that deadly disease, sexually transmitted."(W.Allen)
|
|
|
Rhodanide
Hazard to Others
Posts: 348
Registered: 23-7-2015
Location: Ｔｈｅ ８０ｓ
Member Is Offline
Mood: That retro aesthetic
|
|
| Quote: Originally posted by PHILOU Zrealone | Better start from 1,2-dinitroethan (DNE) and halogenate in basic media to get tetrahalogeno-dinitroethan.
DNE can be obtained via:
-addition of N2O4/NO2 to ethylen (forms also side products nitrite ester, nitrate ester, oxydation products, nitroso, nitro or mix of those)
-addition (substitution) of NaNO2 to 2,3-dihalogeno-butan-1,4-dioic acid (dihalogeno-succinic acid)...mimicking the formation of nitromethane from
chloroacetic acid with NaNO2.
The introduction of a nitro group favourize the elimination of the carboxylic moeity from the carbon holding the NO2 group.
The NO2 may enter as nitrite ester and lead to oxydation products or nitrosation products of the nitro compound (nitronic acids)...urea may reduce the
nitrite ester troubles and side products and increase nitrite salt solubility.
-addition (substitution) of AgNO2 or Na/K/LiNO2 to 1,2-dihalogeno-ethane into a suitable solvent (DMF or DMSO) with urea to reduce the nitrite ester
troubles and side products and increase nitrite salt solubility.
Possible formation of nitro-nitrite ester or dinitrite ester aside from DNE.
[Edited on 4-2-2017 by PHILOU Zrealone] |
Heh heh, uh...
That might be a little much for my limited apparatus, and small basement lab 
|
|
|
![[*]](images/xpblue/default_icon.gif){kind=icon}
