Sciencemadness Discussion Board

Nitronium triflate?

Azane - 4-12-2014 at 20:52

Just an idea that came to me the other day. I don't know if it exists, so I can't definitively say whether it'd be a solid or not (or even stable) at RT or STP, and it seems like the only easy way to produce it in quantitative yields would be to react triflic anhydride with a nitrate source.

However, given the patterns seen in other nitronium salts, few in number may they be, I think it's reasonable to say that NO2SO3CF3 would at least be solid, and it wouldn't be nearly as prone to decomposition as either N2O5 or NO2ClO4. Furthermore, I wouldn't be surprised if it were at least more heat-stable than nitronium salts of Lewis adduct bases (BF4-, SbF6-), and in the case of accidental exposure to water, NO2SO3CF3's hydrolysis wouldn't result in the formation of HF.

I just think it'd be great to have a nitrating agent that wouldn't be an explosion hazard, or a fluoride source, or an unstable liquid solution requiring preparation every time you need to use it. What do you guys think?

deltaH - 4-12-2014 at 21:52

Why do you say it wouldn't be as prone to decomposition as NO2ClO4?

How do you intend to prepare it? When you talk about using the anhydride, were you thinking of this:

N2O5 + (CF3SO2)2O => 2NO2SO3CF3

Bert - 5-12-2014 at 08:13

Quote: Originally posted by Azane  
an explosion hazard


You say that like it's a bad thing?!;)

As you're not directly mentioning explosive/propellant use, and it's just an idea- Let's move this over to Beginnings for now? Glad to move it back with provision of ANY of these: references, experimental data or a detailed proposal for an experiment in production or use of the proposed substance.

Chemosynthesis - 5-12-2014 at 13:40

Organic Chemistry of Explosives page 141 by Jai Prakash Agrawal, Robert Hodgson claims nitronium triflate is probably the agent of nitration in the mixed acid nitration. The citation is given as "G. A. Olah, R. Malhotra and S. C. Narang, Nitration:Methods and Mechanisms, Wiley-VCH,Weinheim (1989)."

Perhaps you can request a copy.

Oh, I found a new article. Gratitude to the USAF:

ANHYDROUS NITRONIUM TRIFLATE NITRATION OF AROMATIC and HETEROAROMATIC COMPOUNDS:CONVENTIONAL BENCHTOP & MICROWAVE-ASSISTED CONDITIONS
by Scott A. Shackelford et al.

pdf warning: http://www.dtic.mil/dtic/tr/fulltext/u2/a440687.pdf

No mention of isolation (generated in situ) or stability, but it appears to be in transition state notation (brackets), so it doesn't appear stable to me.

[Edited on 6-12-2014 by Chemosynthesis]

Chemosynthesis - 6-12-2014 at 13:53

Sorry for the double post, but I can't edit my above post anymore.
I do not see any physical data on the Royal Soc. Chem. chemspider or ... but there is a CAS on ebuychem.com (CAS# 42262-35-1). Unsurprisingly, I didn't get any hits in the Am. Chem. Soc CAS search.
Smiles won't show up with the forum plugin.
SMILES: [CH2+]=[NH+][O-].C(F)(F)(F)S(=O)(=O)[O-]

Seems not to exist in intermediate form as an isolatable substance.

Just as an aside, if it were isolatable, it does appear to be biologically active and bio-available as the deprotonated acid at pH's 7.4 (blood) and 7.2 (intracellular), which likely wouldn't be safe given triflic acid's severe health risks.

chornedsnorkack - 7-12-2014 at 15:26

Quote: Originally posted by Chemosynthesis  


Just as an aside, if it were isolatable, it does appear to be biologically active and bio-available as the deprotonated acid at pH's 7.4 (blood) and 7.2 (intracellular), which likely wouldn't be safe given triflic acid's severe health risks.

Triflic acid is corrosive, yes, but is it toxic?
Regarding the sources of triflyl group: triflyl anhydride is a liquid conveniently boiling at 82 celsius. Triflyl fluoride is a gas boiling at -21 celsius...

How about metathesis with triflates? Like reaction
NO2BF4+CF3SO3K<-> KBF4+CF3SO3NO2... does it by preference go right or left?

DraconicAcid - 7-12-2014 at 15:36

I'd be far more worried about the cation than the triflate anion, personally.

Chemosynthesis - 7-12-2014 at 16:10

Quote: Originally posted by chornedsnorkack  

Triflic acid is corrosive, yes, but is it toxic?
Regarding the sources of triflyl group: triflyl anhydride is a liquid conveniently boiling at 82 celsius. Triflyl fluoride is a gas boiling at -21 celsius...

Hmm. Well, I had read in Efficient Preparations of Fluorine Compounds(ed. Herbert W. Roesky) p48 that triflic acid is extremely toxic, and it's not an old book (2012) however...
Rhodia characterizes it as "not toxic for micro-organisms, not bio-accumulable, not mutagenic,non-sensitizing and has a low accute toxicity" however there are apparently several inconclusive assays on glucocorticoid, estrogen, and mitochondrial membrane disruption. Maybe someone who's worked with it can set me right.

What worries me most healthwise is the extreme acidity, and that I am not sure of the vapor pressure (not listed in the MSDS I have). Rhodia also calls it the "strongest Brønstedt acid available at industrial scale."
Quote: Originally posted by DraconicAcid  
I'd be far more worried about the cation than the triflate anion, personally.

Why would it be any more dangerous than fuming nitric acid?

DraconicAcid - 7-12-2014 at 16:19

Quote: Originally posted by Chemosynthesis  
Why would it be any more dangerous than fuming nitric acid?


It's not. My respect for triflic acid is due to its corrosivity (I have worked with it, and watched it char the Kimwipe that I wiped the syringe off with), not the toxicity of the anion. Methyl triflate is something I've also worked with, and have a great respect for, because I don't want it methylating my DNA.

Chemosynthesis - 7-12-2014 at 16:30

Quote: Originally posted by DraconicAcid  

It's not. My respect for triflic acid is due to its corrosivity (I have worked with it, and watched it char the Kimwipe that I wiped the syringe off with), not the toxicity of the anion. Methyl triflate is something I've also worked with, and have a great respect for, because I don't want it methylating my DNA.

I can appreciate both of those reasons, and your experience. Since I couldn't find a vapor pressure, I am curious; did any of your interactions with triflic acid make inhalation appear to be an issue?

I've been in rooms with some nitrogen oxide plumes venting outside of a hood, and as long you tried not to be next to the fumes or actively breathe it in while fanning it to the nearest hood with something, I did not ever fear for insidious pulmonary edema, which is a health risk.

DraconicAcid - 7-12-2014 at 16:48

It fumed when we opened the Schlenk tube of it, so yes, I'd be very careful of using it without proper ventilation.

Chemosynthesis - 7-12-2014 at 17:05

Very interesting. Thank you for sharing your experience.

chornedsnorkack - 8-12-2014 at 01:41

While I don´t have the saturated vapour pressures at room temperature, let´s compare normal boiling points, where the total vapour pressure equals 1 bar:
Hydrochloric acid, azeotrope (20 %) - 109 degrees
Nitric acid, white fuming (100 %) - 83 degrees
Nitric acid, azeotrope (68 %) - 121 degrees
Sulphuric acid, pure - 280 degrees
Sulphuric acid, azeotrope (98 %) - 337 degrees
Perchloric acid, pure - around 100 degrees, but tends to detonate instead
Perchloric acid, azeotrope (72 %) - 203 degrees
Hydrofluoric acid, azeotrope (36 %) - 111 degrees
Triflic acid, 100 % - 162 degrees

So, triflic acid should be more volatile than sulphuric acid, but less volatile than hydrochloric or nitric acid.
Furthermore, as a strong acid, I should expect triflic acid to form an azeotrope with water boiling higher than the dry triflic acid, but I have not found data of such an azeotrope. Does anyone have them?

DraconicAcid - 8-12-2014 at 09:17

Quote: Originally posted by chornedsnorkack  

Furthermore, as a strong acid, I should expect triflic acid to form an azeotrope with water boiling higher than the dry triflic acid, but I have not found data of such an azeotrope. Does anyone have them?


I only used the anhydrous material. It seems a waste to make such a strong acid and then level it with water.

chornedsnorkack - 8-12-2014 at 10:20

Interestingly, although triflyl anhydride boils around 84 degrees (compare sulphur trioxide 45 celsius), it is hydrated pretty sluggishly, and does not fume.

Chemosynthesis - 8-12-2014 at 10:42

Quote: Originally posted by chornedsnorkack  
While I don´t have the saturated vapour pressures at room temperature, let´s compare normal boiling points: […] So, triflic acid should be more volatile than sulphuric acid, but less volatile than hydrochloric or nitric acid. Furthermore, as a strong acid, I should expect triflic acid to form an azeotrope with water boiling higher than the dry triflic acid, but I have not found data of such an azeotrope. Does anyone have them?
Good reasoning. Sometimes I forget the simplest and most useful estimations in light of the math of Clausius-Clapeyron. Very practical. Speaking of, I should have checked Sigma (due to their nomogram) because they have both vapor pressure and vapor density of triflic acid listed.

I have seen, in literature as well as wikipedia, mentions of purifying triflic acid via distillation with the anhydride, or commercial reagent grade under inert atmosphere prior to use. This would further suggest that an azeotrope does exist.
Ex. citation for inert atm. distillation: PMCID: PMC23573

Perhaps data from DOI: 10.1007/BF02767999 on molar volumes of triflic acid dilutions could be used with an equation similar to one in DOI: 10.1081/LFT-200028058 to estimate an azeotrope boiling point from a liquid activity coefficient?

Azane - 10-12-2014 at 21:22

Quote: Originally posted by deltaH  
Why do you say it wouldn't be as prone to decomposition as NO2ClO4?

How do you intend to prepare it? When you talk about using the anhydride, were you thinking of this:

N2O5 + (CF3SO2)2O => 2NO2SO3CF3


Perchlorates are inherently susceptible to thermal decomposition. Potassium perchlorate decomposes at 400ºC, and all other metal triflates are susceptible to decomposition. Ammonium perchlorate was the main factor in the PEPCON disaster. Nitronium perchlorate itself, being composed of an oxidizing cation and anion, is susceptible to thermal decomposition, as well as detonation. Triflates, on the other hand, are widely known for their stability. I'm not saying that nitronium triflate would be indefinitely stable to heat, just that the triflate would impart favorable properties over anions like perchlorate, nitrate, tetrafluoroborate, and hexafluoroantimonate.

And yes, honestly I also thought of the possibility of using N2O5, but then I realized that Tf2O could be reacted (hypothetically) with any stable nitrate source, such as potassium nitrate, in which case the products would be potassium and nitronium triflate. Again, this is basically just speculation.

[Edited on 11-12-2014 by Azane]

Chemosynthesis - 10-12-2014 at 22:59

Quote: Originally posted by Azane  
And yes, honestly I also thought of the possibility of using N2O5, but then I realized that Tf2O could be reacted (hypothetically) with any stable nitrate source, such as potassium nitrate, in which case the products would be potassium and nitronium triflate. Again, this is basically just speculation.

It has been years since I've touched this and I don't have the software anymore, but you could try to get access to a computational tool such as Gaussian allowing an optimized vibrational frequency (automatically containing thermochemistry analysis) calculation of products and reactants with your basis set of choice.
Use water as your solvent, get the saddle points and minima of the potential energy surface of a reaction. Verify this is a valid transition state using an Intrinsic Reaction Coordinate calculation. Then calculate Gibb's free energies of products, reactants, and proposed transition state.
Now it's less speculative.

chornedsnorkack - 11-12-2014 at 10:07

Even if nitronium triflate does exist, what do you think its melting point is?

Incidentally, can nitronium hydrogen sulphate be separated, and what is its melting point?

deltaH - 11-12-2014 at 11:25

Quote: Originally posted by Azane  
Quote: Originally posted by deltaH  
Why do you say it wouldn't be as prone to decomposition as NO2ClO4?

How do you intend to prepare it? When you talk about using the anhydride, were you thinking of this:

N2O5 + (CF3SO2)2O => 2NO2SO3CF3


Perchlorates are inherently susceptible to thermal decomposition. Potassium perchlorate decomposes at 400ºC, and all other metal triflates are susceptible to decomposition. Ammonium perchlorate was the main factor in the PEPCON disaster. Nitronium perchlorate itself, being composed of an oxidizing cation and anion, is susceptible to thermal decomposition, as well as detonation. Triflates, on the other hand, are widely known for their stability. I'm not saying that nitronium triflate would be indefinitely stable to heat, just that the triflate would impart favorable properties over anions like perchlorate, nitrate, tetrafluoroborate, and hexafluoroantimonate.

And yes, honestly I also thought of the possibility of using N2O5, but then I realized that Tf2O could be reacted (hypothetically) with any stable nitrate source, such as potassium nitrate, in which case the products would be potassium and nitronium triflate. Again, this is basically just speculation.

[Edited on 11-12-2014 by Azane]


Ammonium perchlorate has a cation as a fuel and an anion that is an oxidant, that is why it can explode. Nitronium perchlorate, on the other hand, is all oxidant. It's hypothetical decomposition equation at high temperature during a detonation would be

NO2ClO4(s) => 1/2N2(g) + 3O2(g) + 1/2Cl2(g)

which is nothing else than the opposite of forming it from it's elements, hence the heat of this decomposition at STP would be negative the heat of formation, with dHf being 33.48 kJ/mol [1], hence the heat of decomposition is -33.48kJ/mol. This is not very energetic at all, plus this is the worst case scenario because the nitronium perchlorate at lower temperature would probably decompose to nitric oxides, maybe even nitrosyl chloride and oxygen, for example possibly:

NO2ClO4(s) => NOCl(g) + 5/2O2(g)

This hypothetical decomposition route for this best case scenario is in fact endothermic at 18.23kJ/mol [2].

So it doesn't look that bad to me, unless it's mixed with a fuel that is :) Wiki BTW says that it easily detonable, though I couldn't access the reference to that statement. From my heats calculation, I imagine it would be if you stick a detonator into it, plus the large amounts of oxygen formed would cause a big explosion if their's oxygen deficient fuel products from the detonator.

But aside from purposefully detonating it, I'm not so sure working with it could lead to a detonation easily in the absence of a fuel because of endothermic decomposition products at lower temperature.

References:

[1] Cordes, H.F & Fetter, N.R. The Heat of Formation of Nitronium Perchlorate and of the Gaseous Nitronium Ion. J. Phys. Chem., 1958, 62 (10), pp 1340–1341.

[2] Using dHf for NOCl(g) = 51.71kJ/mol from the NIST webbook, http://webbook.nist.gov/cgi/cbook.cgi?ID=C2696926&Units=...

***
The reason I suggested preparing it from N2O5 was not to waste triflate as it's expensive. The hypothetical N2O5 route plus triflic anhydride is also potentially cleaner, producing no potassium triflate salt to have to separate, though preparing N2O5 is a harder route.

[Edited on 12-12-2014 by deltaH]

Chemosynthesis - 11-12-2014 at 12:11

Quote: Originally posted by chornedsnorkack  
Even if nitronium triflate does exist, what do you think its melting point is?

Incidentally, can nitronium hydrogen sulphate be separated, and what is its melting point?

Easy enough to plug them into ChemBioDraw and get a quirky extrapolated QSPR value I wouldn't trust.

chornedsnorkack - 12-12-2014 at 05:02

Quote: Originally posted by deltaH  
Quote: Originally posted by Azane  
Quote: Originally posted by deltaH  
Why do you say it wouldn't be as prone to decomposition as NO2ClO4?

How do you intend to prepare it? When you talk about using the anhydride, were you thinking of this:

N2O5 + (CF3SO2)2O => 2NO2SO3CF3


Perchlorates are inherently susceptible to thermal decomposition. Potassium perchlorate decomposes at 400ºC, and all other metal triflates are susceptible to decomposition. Ammonium perchlorate was the main factor in the PEPCON disaster. Nitronium perchlorate itself, being composed of an oxidizing cation and anion, is susceptible to thermal decomposition, as well as detonation. Triflates, on the other hand, are widely known for their stability. I'm not saying that nitronium triflate would be indefinitely stable to heat, just that the triflate would impart favorable properties over anions like perchlorate, nitrate, tetrafluoroborate, and hexafluoroantimonate.


Ammonium perchlorate has a cation as a fuel and an anion that is an oxidant, that is why it can explode. Nitronium perchlorate, on the other hand, is all oxidant. It's hypothetical decomposition equation at high temperature during a detonation would be

NO2ClO4(s) => 1/2N2(g) + 3O2(g) + 1/2Cl2(g)

which is nothing else than the opposite of forming it from it's elements, hence the heat of this decomposition at STP would be negative the heat of formation, with dHf being 33.48 kJ/mol [1], hence the heat of decomposition is -33.48kJ/mol. This is not very energetic at all, plus this is the worst case scenario because the nitronium perchlorate at lower temperature would probably decompose to nitric oxides, maybe even nitrosyl chloride and oxygen, for example possibly:

NO2ClO4(s) => NOCl(g) + 5/2O2(g)

This hypothetical decomposition route for this best case scenario is in fact endothermic at 18.23kJ/mol [2].

So it doesn't look that bad to me, unless it's mixed with a fuel that is :) Wiki BTW says that it easily detonable, though I couldn't access the reference to that statement. From my heats calculation, I imagine it would be if you stick a detonator into it, plus the large amounts of oxygen formed would cause a big explosion if their's oxygen deficient fuel products from the detonator.


Look at it this way: nitronium triflate and nitronium perchlorate are effectively mixed anhydrides.

Nitronium perchlorate is a mixed anhydride of N2O5 and Cl2O7.
N2O5 readily decays
2N2O5->4NO2+O2
Mind you, N2O5 is not heard to explode without a reducer. Probably that decay process has small enthalpy but a low kinetic barrier. Still, it goes easily.
Cl2O7 readily detonates releasing a lot of heat:
2Cl2O7->2Cl2+7O2
This reaction is kinetically harder than N2O5 decay, which is why it is explosive.
If you have a mixture of N2O5 and Cl2O7 then when Cl2O7 explodes, it disintegrates N2O5 as well
N2O5+Cl2O7->2NO2+Cl2+4O2
NO2 may dilute some heat from Cl2O7 decay, but not enough to prevent a thermal runaway.
Now, nitronium perchlorate... How much energy is released by the reaction
N2O5+Cl2O7->2NO2ClO4?

And now compare nitronium triflate. Triflic anhydride is
(CF3SO2)2O
It is resistant to decay, and most decay paths take strong heating.
How much energy is released by reaction
N2O5+(CF3SO2)2O<->2NO2OSO2CF3
How well does triflate stabilize nitronium, and what forms when that triflate is finally pyrolysed?