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Author: Subject: Nitroalkane syntheses ?
AndersHoveland
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[*] posted on 5-3-2013 at 22:36


Quote: Originally posted by Formatik  

The reaction is not at room temperature (60 C), though the original reference basically only showed this to work in cycloalkanes, and one longer chain alkane.

I really do not think it matters too much. The reaction is very versatile.

Here is an informative link:
http://www.tandfonline.com/doi/pdf/10.1080/01614940902743841


Quote: Originally posted by zed  
Well ethylene is widely available. Addition of HNO2, across the double-bond might work in some universe.

We know that nitrogen dioxide can react with ethylene under ambient conditions to form 1,2-dinitroethane. It procedes through a radical mechanism. (if there is any air some mixed nitrate esters also form)

Perhaps if the reaction was done in solution in the presence of Fe+2 ions. Using an excess of ethylene in solution, and slowly pass NO2 in, while also periodically making small additions of an Fe+2 compound in intervals.

Normally ferrous salts reduce NO2 to NO, but if some other inert solvent (besides water) was used I would imagine that FeNO2+2 would form. I do not know, but possibly if this substance was heated in the presence of ethylene. Many nitrites tend to be rather unstable; solutions of ammonium nitrite, for example, decompose even before the boiling point of water is reached.

The radical •CH2-CH2NO2 would probably tautomerize to •CH2-CH=NO2-, and would be able to immediately oxidize the Fe+2 ion (also in the presence of two hydrogen ions) to finally form nitroethane.

Just an idea.

[Edited on 6-3-2013 by AndersHoveland]
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[*] posted on 6-3-2013 at 07:24


Quote: Originally posted by AndersHoveland  
The reaction is very versatile.

Only if the substrate is within its scope. Methane and ethane do not appear as such since neither has a reactive methine, methylene or benzylic group that could be subjected to N-hydroxyphthalimide catalyzed nitration. It is apparent from the examples and the reaction regioselectivity, as described in that article, that methyl groups do not undergo the reaction. That is something one can infer also from the proposed reaction mechanism.
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[*] posted on 6-3-2013 at 18:45


That being said, what are the possibilities of producing Nitroethane, via Ethylene?
I haven't seen any literature on the topic, but it does seem like it might be possible.

[Edited on 7-3-2013 by zed]

[Edited on 7-3-2013 by zed]
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[*] posted on 6-3-2013 at 20:42


Quote: Originally posted by Nicodem  

Only if the substrate is within its scope. Methane and ethane do not appear as such since neither has a reactive methylene group

You really think that a methyl group would be less reactive than a methylene group?! Please show me just one other example in organic chemistry where this is the case (where it is only bonded to other carbon atoms). Ethane is an alkane just like propane or cyclohexane.

[Edited on 7-3-2013 by AndersHoveland]
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[*] posted on 6-3-2013 at 23:05


N-hydroxyphthalimide catalysis: the first paper mentioned about this wasn't clear which substrates would work. But, looking here: http://dx.doi.org/10.1021/jo025632d it apparently does work for lighter alkanes (CH3CH3, CH3CH2CH3), though at higher temperatures (100 C for 14 hrs, working at 50 C produced a miserable yield (6%) i.e. in nitropropane).

This would all need an autoclave of course.


The older literature mentions ethylene nitrite C2H4(NO2)2, alongside a volatile poisonous oil, being formed from passing ethylene into liquid N2O4, or heating the two gases (Dictionary of Chemistry by H. Watts).
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[*] posted on 17-3-2013 at 17:16
Nitroethane from 2-Bromoproprionic acid


I'm not sure if I should have started a new thread or not. If so, I apologize.

Of the many syntheses of nitroethane discussed in this thread, one method in particular sparked my interest. UpNatom, in his posts above, mentioned the method described in US patent 4,319,059, using 2-Bromoproprionic acid, Rhodium started a thread about this patent on the Hive too. The patent describes a straightforward reaction of 2-bromoproprionic acid with an alkali nitrite and magnesium ion (methoxide is preferred but the sulfate, chloride, etc. are also said to work) in DMSO. Since 2-bromoproprionic acid can be made from alanine using HBr and sodium nitrite, I thought I'd give it a try. Here are the results:

Experimental 1: 85 g of magnesium sulfate was added to 200 mL of DMSO and heated on a hot-water bath to help the MgSO4 dissolve, which was a waste of time because hardly any dissolved. 50 g of 2-bromoproprionic acid was poured into the flask and swirled as magnetic stirring didn't work. An attempt was made to dissolve 28 g of sodium nitrite in DMSO but it didn't dissolve well either. Maybe more time was needed. The sodium nitrite/DMSO mixture was poured into the flask quickly resulting in a slightly exothermic reaction. It may be better to add the NaNO2 as powder in increments. The flask was stoppered and manually shaken once in awhile over about 5 hours. The pressure was released by removing the stopper occassionally. It popped out a few times. Mechanical stirring and a reflux column would be better for this reaction. I ended up allowing the mixture to sit overnight, which isn't a problem according to the patent. In the morning the mixture was poured into a 1000 mL sep funnel, filtering the still undissolved magnesium sulfate, added a couple of splashes of 14 HCl and nearly filled the remaining space in the sep funnel with water. I extracted this mixture with DCM (4 x 100 mL), washed the extracts with brine, and dried over MgSO4, before removing the DCM by vacuum. The residue consisted of what appeared to be stinky, hot DMSO. No nitroethane was recovered.

Experimental 2: 430 mL of DMSO was poured into a 2000 mL two-neck flask equipped with mech. stirring and a reflux column. The column was removed while making the additions. 30 g of magnesium methoxide was added and stirred for 30 min. 50 g of 2-bromoproprionic acid was poured into the flask and allowed to stir for 15 min. before adding 25 g of sodium nitrite in 5 g increments over an hour. The mixture was stirred at room temperature for 5 hours. 600 mL of 10% HCl was added carefully to neutralize the reaction; this was stirred for 30 min. Extraction with DCM (4 x 100 mL). I skiipped washing the extracts and went straight to distillation, removing the DCM and then vacuum distilling the stinky residue. No nitroethane was collected.

I can smell a sweet smell that I assume is nitroethane after the reaction is complete. The water layer smelled more like nitroethane than the DCM extract. It doesn't appear that the nitroethane is going into the DCM layer but I'm not entirely sure there is nitroethane. My wife ended up dumping the aqueous layer down the drain. This is my first attempt to use DMSO and I must tell you, it stinks.

Any help would be appreciated.
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[*] posted on 18-3-2013 at 23:56


Quote: Originally posted by anomolous  
Of the many syntheses of nitroethane discussed in this thread, one method in particular sparked my interest. UpNatom, in his posts above, mentioned the method described in US patent 4,319,059, using 2-Bromoproprionic acid, Rhodium started a thread about this patent on the Hive too. The patent describes a straightforward reaction of 2-bromoproprionic acid with an alkali nitrite and magnesium ion (methoxide is preferred but the sulfate, chloride, etc. are also said to work) in DMSO. Since 2-bromoproprionic acid can be made from alanine using HBr and sodium nitrite, I thought I'd give it a try.


What if you did this reaction on glutamic acid (from the common food additive MSG) ? Presumably you could get 4-nitro-butyric acid.
NO2-CH2CH2CH2-COOH

Now, does anyone remember the thread about making nitromethane from 2-chloroacetic acid? I am not going to go into the details, but will quickly summarize.

One of the possible ways to chlorinate acetic acid is to dissolve it in acetic anhydride and pass chlorine into the solution. Under the extremely acidic conditions, acetic acid essentially reverts to its CH2=C(OH)2 tautomer, through which it can be attacked by the chlorine.

Presumably, we could selectively chlorinate 4-nitro-butyric acid under the same conditions to get 2-chloro-4-nitro-butyric acid.
NO2-CH2-CH2-CHCl-COOH

Nitromethane can be made by distilling 2-chloroacetic acid with sodium nitrite. So what would happen if we distilled 4-nitro-butyric acid with sodium nitrite? Is it possible that 1,3-dinitropropane could result?
NO2-CH2-CH2-CH2-NO2

The thing about this compound is that the nitro groups are not vicinal, so the compound would not display thermal instability. (it would not be a good idea to try to distill 1,2-dinitroethane for example)

[Edited on 19-3-2013 by AndersHoveland]




I'm not saying let's go kill all the stupid people...I'm just saying lets remove all the warning labels and let the problem sort itself out.
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[*] posted on 20-3-2013 at 22:44


The reason I was attracted to this route was its apparent simplicity. I may try vacuum distilling instead of extracting with a solvent. But I don't understand why a solvent extraction wont work? Has anyone dealt with DMSO before? It seems to have strange characteristics. Is there something I'm missing? Do you think vacuum distillation is the answer or is there some other method of extracting the nitroethane you'd recommend?

Another method using 2-bromoproprionic acid was mentioned on the Rhodium archive. Nitroethane Synthesis: A Compilation, method 7. It is the destructive distillation of 2-bromoproprionic acid. It occurs in a solution of K2CO3 with NaNO2. Follows is the reference: V. Auger. Bull Soc. Chim. France Post no 3,23,333 (1900). If you can find it please share. Though the yields are so-so (50%) I may give this method a try. But the patent method still haunts me.

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[*] posted on 24-3-2013 at 08:18


Another semi-failed attempt:

10 g magnesium methoxide in 50 mL of DMSO, followed by 17 g 2-bromoproprionic acid. 9 g NaNO2 dissolved in 65 ml DMSO added dropwise then stirred for 12 hours. Neutralized with 15 mL of 31% HCl. Vacuum distilled collected about 15 mL of liquid. Scent of DMSO present and faint fruity scent present. Added about 25 mL of water then extracted with DCM (2 x 10 mL). The DCM evaporated leaving a few grams of nitroethane ( I think).

Questions: How can I test for nitroethane? Do you think the 2-bromoproprionic acid should be pure? Would ether be a better solvent? Why is no one interested in this reaction?
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[*] posted on 24-3-2013 at 10:45


Quote: Originally posted by anomolous  
If you can find it please share.


This thread?




"You're going to be all right, kid...Everything's under control." Yossarian, to Snowden
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[*] posted on 26-3-2013 at 19:03


Quote: Originally posted by AndersHoveland  
Quote: Originally posted by anomolous  
Of the many syntheses of nitroethane discussed in this thread, one method in particular sparked my interest. UpNatom, in his posts above, mentioned the method described in US patent 4,319,059, using 2-Bromoproprionic acid, Rhodium started a thread about this patent on the Hive too. The patent describes a straightforward reaction of 2-bromoproprionic acid with an alkali nitrite and magnesium ion (methoxide is preferred but the sulfate, chloride, etc. are also said to work) in DMSO. Since 2-bromoproprionic acid can be made from alanine using HBr and sodium nitrite, I thought I'd give it a try.


What if you did this reaction on glutamic acid (from the common food additive MSG) ? Presumably you could get 4-nitro-butyric acid.
NO2-CH2CH2CH2-COOH

Now, does anyone remember the thread about making nitromethane from 2-chloroacetic acid? I am not going to go into the details, but will quickly summarize.

One of the possible ways to chlorinate acetic acid is to dissolve it in acetic anhydride and pass chlorine into the solution. Under the extremely acidic conditions, acetic acid essentially reverts to its CH2=C(OH)2 tautomer, through which it can be attacked by the chlorine.

Presumably, we could selectively chlorinate 4-nitro-butyric acid under the same conditions to get 2-chloro-4-nitro-butyric acid.
NO2-CH2-CH2-CHCl-COOH

Nitromethane can be made by distilling 2-chloroacetic acid with sodium nitrite. So what would happen if we distilled 4-nitro-butyric acid with sodium nitrite? Is it possible that 1,3-dinitropropane could result?
NO2-CH2-CH2-CH2-NO2

The thing about this compound is that the nitro groups are not vicinal, so the compound would not display thermal instability. (it would not be a good idea to try to distill 1,2-dinitroethane for example)


How are you going to make 4-nitrobutyric acid from glutamic acid?

Acetic acid and acetic anhydride are not what I would consider "extremely acidic" conditions. You have the mechanism wrong.

You don't know that you can selectively chlorinate 4-nitrobutyric acid at the 2 position. You're just guessing.

http://pubs.acs.org/doi/abs/10.1021/ja01170a039

How do you know that 1,3-dinitropropane is thermally stable enough to distill at any pressure up to 760mm Hg? (Implied by your failure to specify reduced pressure.) You don't, and again, you're guessing. What if you're wrong, and someone tries to distill it and it detonates? What then? Has it ever occurred to you that what you post could kill someone?

Quote:
You really think that a methyl group would be less reactive than a methylene group?! Please show me just one other example in organic chemistry where this is the case (where it is only bonded to other carbon atoms). Ethane is an alkane just like propane or cyclohexane.


Organic chemistry is full of subtlety. Carbon substituted aliphatic carbons are activated compared to unsubstituted ones. For example, look at the rate of peroxide formation with isopropylbenzene vs. ethylbenzene. You cannot just assume that what will work on one system will work on an apparently similar one, as it often will not.

You need to spend much more time studying fundamentals and a lot less scurrying from one random paper to the next, accumulating vast tomes of information that is ultimately useless to you as you are unable to integrate it. Some time spent with March's Advanced Organic Chemistry, and Physical Organic Chemistry by Anslyn and Dougherty, would do you a lot of good.

[Edited on 27-3-2013 by madscientist]




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[*] posted on 27-3-2013 at 09:29


Quote:
How are you going to make 4-nitrobutyric acid from glutamic acid?


I should clarify that I understand your plan is to swap out the amino group for a bromo and then a nitro, and decarboxylate. But how are you going to isolate it? I assume chromatography is out. That doesn't leave you with a lot of good options.

Impurities are like a plague, they multiply exponentially every time you carry a crude over to the next step. One pot procedures require careful consideration to develop.

The workup is usually half the battle, and sometimes a synthetic pathway will be deemed unfeasible due to a single difficult isolation. Practicality is something you need to be thinking about more often than never.




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[*] posted on 27-3-2013 at 15:19


Destructive Distillation of 2-bromoproprionic acid

120 mL of K2CO3 sol. (pH 11) was poured into a 500 mL flask, followed by 40 g of crude 2-bromoproprionic acid. Then 40 g of NaNO2 was added carefully in 5 g increments. After the NaNO2 had dissolved the reaction flask was placed into a preheated oil-bath (145C) and the distillation begun. The distillate separated into two layers with nitroethane on the bottom. The nitroethane was extracted with ether and ether removed via distillation, yielding 16 g nitroethane.
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[*] posted on 30-3-2013 at 21:04


I tried the reaction of 2-bromoproprionic acid with sodium nitrite in DMSO again. This time I vacuum distilled the resulting mixture yielding only a tiny bit of nitroethane. Is this patent bogus? I've gotten better results with the destructive distillation of 2-bromoproprionic acid. Bummer. Feel free to chim in if you learn anything new.
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[*] posted on 31-3-2013 at 19:43


Also to correct a mistake, obviously I meant 2-chloro-4-nitro-butyric acid being distilled with sodium nitrite.

Quote: Originally posted by madscientist  
You don't know that you can selectively chlorinate 4-nitrobutyric acid at the 2 position. You're just guessing.

Do a search using the words "α-chlorination of carboxylic acids".

Quote: Originally posted by madscientist  

http://pubs.acs.org/doi/abs/10.1021/ja01170a039

How do you know that 1,3-dinitropropane is thermally stable enough to distill at any pressure up to 760mm Hg? (Implied by your failure to specify reduced pressure.) You don't, and again, you're guessing. What if you're wrong, and someone tries to distill it and it detonates? What then? Has it ever occurred to you that what you post could kill someone?

I think you are just getting outrageous here. I cannot imagine it is any more dangerous than nitromethane. If you think about it, it is just two nitromethanes bonded to a methylene group. Even your link mentioned nothing about any explosion danger. Just to allay any misplaced fears, a quick search revealed the boiling boint of 1,3-dinitropropane is 263.6 °C at 760 mmHg. Distillation is probably not even a practical option. Separation should still be simple enough since it is oil soluble and would likely separate out by itself.

Quote: Originally posted by madscientist  
Carbon substituted aliphatic carbons are activated compared to unsubstituted ones. For example, look at the rate of peroxide formation with isopropylbenzene vs. ethylbenzene.

Okay, yes I know. But that is a carbon bonded to three other carbons and only one hydrogen. I do not suppose you know of any examples where a carbon with two hydrogens and two carbons is more reactive? The example you described is only more reactive because the lone C-H bond is so much more polar.

Quote: Originally posted by madscientist  
But how are you going to isolate it? I assume chromatography is out. That doesn't leave you with a lot of good options.

Impurities are like a plague, they multiply exponentially every time you carry a crude over to the next step.

Your objections are noted. The only potential troublesome impurity that I can foresee would be secondary amines left after the initial diazotization/bromination steps.

I realize my proposal is a long route, but for amateur chemists it still might be easier than any of the alternatives. 1,3-dinitropropane is not an easy compound to synthesize from common chemicals.

[Edited on 1-4-2013 by AndersHoveland]
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[*] posted on 1-4-2013 at 08:10


http://msdssearch.dow.com/PublishedLiteratureDOWCOM/dh_07c8/...

Quote:
Elevated temperatures can also increase ease of detonation; do not distill nitromethane.


Nitromethane boils around 100C. If 1,3-dinitropropane boils at 260C (what's the source of this data?), then that is incomparably more dangerous. And what you're failing to understand - more proof you have no actual lab experience - is that boiling points can be determined even if something decomposes or poses a detonation hazard at those temperatures.

Quote:
Distillation is probably not even a practical option.


You say that now. But you previously implied otherwise:

Quote:
The thing about this compound is that the nitro groups are not vicinal, so the compound would not display thermal instability. (it would not be a good idea to try to distill 1,2-dinitroethane for example)


More wisdom:

Quote:
I think you are just getting outrageous here. I cannot imagine it is any more dangerous than nitromethane. If you think about it, it is just two nitromethanes bonded to a methylene group.


Yes, and by this same logic, there's no reason why the hazards of methyl nitrate, ethylene glycol dinitrate, and nitroglycerine would differ in any way. Long story short, you are guessing as always - assuming that things are as simple as you want to believe they are, and that the results will be as you hope they are.

Quote:
Even your link mentioned nothing about any explosion danger.


Journal articles usually don't mention the hazards of the chemicals they're working with unless they are truly exceptional. They assume that their readers aren't stupid enough to heat an explosive to 260C and let it bump away in a giant RB flask for an hour or two. Perhaps they are mistaken.

Quote:
Okay, yes I know. But that is a carbon bonded to three other carbons and only one hydrogen. I do not suppose you know of any examples where a carbon with two hydrogens and two carbons is more reactive? The example you described is only more reactive because the lone C-H bond is so much more polar.


The yields of 1-bromopropane and 2-bromopropane are not the same. At 150°, the relative yield
of 1-bromopropane is 8% of the monobrominated products, while 2-bromopropane has a 92%
relative yield. This result is surprising since if all H's on propane were equally reactive to Br.,
the relative yield of1-bromopropane should be 75%, and that of 2-bromopropane should be 25%.

Source: http://web.chem.ucsb.edu/~neuman/orgchembyneuman.book/11%20R...


This means bromination of a methylene in this system is nearly 40x faster than that of a methyl. There's a big difference and your insistence to the contrary really is just proof that you have zero credibility.

Quote:
Do a search using the words "α-chlorination of carboxylic acids".


I know about the Hell-Volhard-Felinskii reaction. Here's the problem:

Juvell has studied the velocity of bromination and chlorination of nitromethane, mono- and di-bromonitromethane and secondary nitropropane in 1N hydrogen halide. He finds that the reaction is always of the first order with respect to the nitrocompound. The velocities of bromination and chlorination are the same, and the concentration of the halogen is without influence. Consequently, it must be assumed that the velocity is determined by the unimolecular rearrangement of the nitrocompound to the aci-form, which takes up halogen instantaneously.

Source: http://www.sdu.dk/media/bibpdf/Bind%2010-19%5CBind%5Cmfm-12-...


When you are considering a new synthesis, you have to look at every possible route that the chemistry could take, not just the one you want to happen. And you need to prove, not hope, that these other routes won't predominate, or do a small scale test in the lab to find out (small means under 100mg, bigger than that and unexpected results could become major hazards).

Your post in another thread, advising people to run a SO<sub>3</sub> reaction in CCl<sub>4</sub> was evidence of your incompetence or laziness in this regard, and could have (and perhaps already has) cost lives, as this is, after all, how the Italians made their phosgene in WWI.

Quote:
The only potential troublesome impurity that I can foresee...


Your foresight is not of any value, you have zero credibility whatsoever, you've mostly just cluttered up the forum with bad (often lethal) advice, and speculation with no references, or references that do not back up your point. It would be much preferred if you were to simply keep your great ideas to yourself for now.




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[*] posted on 1-4-2013 at 11:25


Quote: Originally posted by AndersHoveland  
The only potential troublesome impurity that I can foresee would be secondary amines left after the initial diazotization/bromination steps.

I generally don't have the time and patience to deal with the immense stupidity that Anders spams all over the forum and can only admire madscientist's efforts, but this sentence sounds like a chemistry April's fool joke. No chemist with a minimum of synthesis experience would be so narrow sighted to ever say something so chemically oxymoronic.

Anders, you will never be able to understand why your ideas are the result of incompetence, because for being able to comprehend that, you would need at least some competence, which you clearly don't have any. Your recent signature is clear evidence of this disability of which you would be the first victim, if you would ever become honest: "I'm not saying let's go kill all the stupid people...I'm just saying lets remove all the warning labels and let the problem sort itself out."




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[*] posted on 2-4-2013 at 11:37


So this is how information gets buried under tons of misinformation and unrelated dialog. Does anyone want to comment on my post? I didn't think so.
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[*] posted on 2-4-2013 at 11:40


Experimental 1: 300 mL of 96% H2SO4 was carefully added to 600 mL of ~95% ethyl alcohol in a 1000 mL Erlenmeyer. 250 g of anhy. NaSO4 was added and the mixture placed into the fridge and then into the freezer to cool to about 10C. The hydrated sodium sulfate was filtered leaving ~800 mL of ethylsulfate/alcohol solution . The solution was treated with sodium carbonate until a neutral pH of 7 was reached. Blue indicator paper turned slightly pink and red paper turned slightly blue at this pH. The neutral solution was cooled to near 0C and the precipated salt (mostly sodium sulfate) was filtered. The filtered solution was distilled and about 200 mL of alcohol was collected. The residue was filtered again and evaporated in a pyrex dish until a skin began to form, at which time the heat was turned off and the solution allowed to cool on the hotplate. A lot of salt, consisting of a mixture of sodium sulfate, carbonate, and ethyl sulfate, precipitated. This salt was saved to be analyzed and possibly used to produce more nitroethane. The filtered liquid, about 292g/240 mL, and a density of 1.21 g/mL was further distilled and more salts filtered yielding 210 g/155 mL (= Density of 1.35 g/mL). This close to the density of a 50-55% solution of ethyl sulfate, according to the Antonch and Desseigne figures. Upon sitting further, overnight, the solution crystalized further, leaving 164 g/123 mL (Density = 1.33) of sodium ethylsulfate sol. A 55% (by weight) solution of 210 g ethylsulfate sulfate sol. would contain 210 g * .55 = 115.5 g. The molar mass of ethyl sulfate is 148.12, which means we hypothetically had about 0.77 moles of ethyl sulfate in solution. I will react this solution with a molar ratio of sodium nitrate—about 53 g. 7 g of potassium carbonate (~6.9% based on 210 g sodium ethylsulfate sol.) will be used and 6 g of cetyl alcohol hair-conditioner. e
6 g of hair-conditioner was added to a 500 mL flask equipped for distillation, followed by 7g of potassium carbonate dissolved into 27 mL of water. Stirring is activated and 53 g of sodium nitrite is added. The oil bath is heated to 155C and the sodium ethylsulfate sol is added dropwise. The addition was stopped because it took some time for distillate to occur then continued at the rate of distillation. After the addition 100 mL of water was added dropwise and the distillation continued to make sure all of the nitroethane distilled over. About 150 mL of distillate, a mixture of water and yellow globules of nitroethane, was collected. This was extracted with ether (3 x 50 mL). The ether was removed by distillation, and the residue, maybe 15 mL, was saved to be further purified once more is collected.

Experimental 2: 600 mL of 95% ethyl alcohol was carefully mixed with 300 mL of 96% sulfuric acid. A titration was done showing it was a 9.26 molar solution. 840 mL required 7.78 mole Ca(OH)2 is required to neutralize this solution so 575 g of Ca(OH)2 was mixed with 2000 mL of water in a large pitcher and the 840 mL of ethylsulfate/alcohol solution was carefully added to this with stirring. After the the ethylsulfate solution had been reacted with the Ca(OH)2 another 500 mL of water was added to assist in extraction and filtering as the mixture was very thick. The CaSO4 formed was filtered through clothe and then through filters until the solution was clear, resulting in 1970 mL of CaEtHSO4 solution. 100 mL of this solution was reacted with an excess of Na2CO3 and the chalk vacuum filtered, dried in an oven, and weighed (4.4 g) to determine the amount of Na2CO3 needed for neutralization. In this case 0.87 mol or 91 g of Na2CO3 (4.4 g/100 = 0.044 mol; 0.044 mol x 19.7 = 0.87 mol; 0.87 * 106 = 91 g). I used 92 g as an excess was recommended. The resulting solution had a pH of 9.4 and turned red indicator paper a vibrant blue. The mixture was vacuum filtered. Do not try to filter all of the CaCO3 all at once unless you have a gigantic filter because it will not fit. I had a fairly large Buchner and did half of the mixture; I rinsed the chalk, emptied the Buchner funnel, changed the filter, and then did the other half. The filtrate was refiltered by gravity filtration until clear and then evaporated on a pyrex cakepan with a fan blowing over it overnight. CaCO3/NaEtHSO4 salts precipitated during the evaporation. These were filtered and saved for further experimentation and the filtrate, 256 g/200 mL (Density: 1.28), was used in the next step. It will be assumed that 0.87 mol or 124 g of sodium ethyl sulfate is in solution and this will be reacted in a molar ratio to sodium nitrate (58 g). To a 500 mL flask with three-neck adapter, equipped with mechanical stirring, for distillation, and with an addition funnel, was added a squirt (5-7 g) of cetyl alcohol hair-conditioner, 30 mL of water, and 8 g of potassium carbonate. This was mixed, placed into an oil-bath and the oil-bath heated quickly to 155C with occasional stirring. Once the oilbath reached 155C the addition funnel was charged with the NaEtHSO4 solution and the addition began. The addition was carried out very slowly at first and then accidently nearly all of the NaEtHSO4 was added within 20 min.I allowed this mixture to distill for a long time before adding the last 20-30 mL. The distillate, water with yellow oil globules present was extracted with ether and the ether removed by distillation. The residue (12 mL) was saved to be purified further when more nitroethane has been collected.

Experimental 3: The sodium ethylsulfate sol. was prepared as above, except the titrations and sample reactions were not done; the ethylsulfate sol. made by reacting 600 mL 95% ethyl alcohol and 300 mL 96% H2SO4 was reacted with amounts as above (i.e. 575 g Ca(OH)2, 92 g Na2CO3) and then evaporated to a manageable amount (125 mL). The NaEtHSO4 solution was reacted with 60 g of NaNO2 in 30 mL of H2O with 10 g of K2CO3 and a squirt of cetyl alcohol hair conditioner. After the addition 50 mL of water was added. About 400 mL of distillate was collected. The yellow globules of nitroethane were apparent. This was extracted with ether (2 x 50 mL). The ether removed by distillation and the residue (22 mL) was added to the previously collected nitroethane.
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organichem
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[*] posted on 20-11-2013 at 07:25


A maybe interesting article I found:
Improved chemoselective, ecofriendly conditions for the conversion of primary alkyl halides into nitroalkanes under PEG400
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[*] posted on 8-12-2013 at 17:42


Quote: Originally posted by organichem  
A maybe interesting article I found:
Improved chemoselective, ecofriendly conditions for the conversion of primary alkyl halides into nitroalkanes under PEG400

Thank you for the article link!

Attachment: Improved chemoselective, ecofriendly conditions for the conversion of primary alkyl halides into nitroalkanes under PEG4 (87kB)
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[*] posted on 25-12-2013 at 09:29


@Formula409,
Thanks for the attached article.

Very interesting...
Maybe it could work with lower MW PEG like diethylene glycol, triethylene glycol, ...
Should also work with 1,4-dioxane (O(-CH2-CH2-)2O)... or related crown ethers and cryptants...





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[*] posted on 25-12-2013 at 09:43


I hope the Epsom Salts were anhydrous



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[*] posted on 21-2-2017 at 17:28


If i try to do nitration of butane with NO2 in regular steel pipes can they sustain that for just one time?
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[*] posted on 22-2-2017 at 11:39


What about ethyl iodide and sodium nitrite under high pressure?
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