If you is considering partaking in N2O infused beverages, available both in some bars and at home, you make wish to make sure that you avoid letting
your new favorite drink enjoy too much sunlight, especially if you are working with the 95% mixing grain alcohol.
Here is some background: N2O apparently is much more soluble in alcohol, vegetable oil, ...than in water. This can be used to effect chemical reaction
in the lab and at home to infuse flavorings into beverages (link http://www.cookingissues.com/2010/08/11/infusion-profusion-g... . To quote:
"You can infuse flavors into liquor (and water based things, too) almost instantly with nothing more than an iSi Cream Whipper . You can use seeds,
herbs, spiced, fruits, cocoa nibs, etc."
The possible explanation per a comment by Dan, to quote:
"Well done! You’ve rediscovered nitrogen cavitation and put it to novel use. Cavitation is used in cell biology labs to gently disrupt cells in a
vessel ominously called a “nitrogen bomb” (as the gas here is nitrogen). The usual explanation is close to what you surmised: under pressure, gas
penetrates the cells, then forms bubbles to disrupt the cells when the pressure is suddenly released. Your use of the technique looks a lot more
interesting than my experiences with it in a lab."
Now, per the chemistry I outline below, the presence of any formed Ethyl nitrite (possibly formed from the photolysis of the N2O infused Ethanol) may
be unpleasant, to say the least. It turns out that C2H5NO2 is a highly toxic substance capable of producing significant headaches and in higher doses,
even mental illness and death.
"Another way that N2O is lost is via photolysis. An energetic UV photon is able to dissociate N2O into molecular nitrogen, N2, and the singlet D
oxygen atom.
N2O+hc/lambda --> N2+O(1D) "
Then, having created this active form of oxygen, per the same source, reactive nitrogen species are formed from N2O via the reaction of dinitrogen
oxide with the singlet D oxygen atom to form two molecules of nitric oxide as follows:
N2O + O(1D) --> 2 NO
which in the presence of air/O2, can form NO2:
2 NO + O2 --> 2 NO2
The problem is the formation of NO and NO2, as per this source: "Spirit of nitrous ether and ethyl nitrite: a bibliographical index of ...", by W. O.
Richtmann, J. A. Anderson, available at http://books.google.com.bz/books?pg=PA34&id=7yfrAAAAMAAJ... to quote:
"Kopp prepares ethyl nitrite by using a mixture of alcohol, nitric acid, copper turnings, and passing the gas through a solution of ferrous sulphate
to remove nitrous acid and other oxides of nitrogen. Gerhardt. 1851."
As the action of copper on Nitric acid will produce either NO (or NO2 with concentrated HNO3), and as the action of UV light on N2O with activated
oxygen can apparently also as discussed above, this appears to be an alternate preparation of the highly toxic Ethyl nitrite. See MSDS at https://www.google.com/url?sa=t&source=web&rct=j&... A most ironic thing about the possible formation of this compound, as per the
cited MSDS, it smells like rum. My take on the chemistry:
2 C2H5OH + NO + NO2 = 2 C2H5N02 + H2O
So, an excess of water (working with lower proof alcohol) would be safer than so called mixing alcohol and certainly avoid excess of sunlight when
trying to indulge with your N2O infused beverage (like at the beach), as otherwise you may develop quite a kicker (even with just low doses of any
Ethyl nitrite formed).
[Edited on 9-8-2014 by AJKOER]unionised - 9-8-2014 at 06:15
In the real world nitrous oxide only absorbs UV with wavelengths shorter than 200 nm.*
Since oxygen also absorbs those wavelengths strongly there is no such radiation in sunlight at ground level.
The photochemical reaction AJKOER cites simply won't happen.
* for those who are interested, it has absorption bands near 118 nm, 129 nm, 147 nm and 185 nm.
hissingnoise - 9-8-2014 at 12:08
Quote:
Another way that N2O is lost is via photolysis. An energetic UV photon is able to dissociate N2O into molecular nitrogen, N2, and the singlet D oxygen
atom.
Atomic oxygen is not "singlet oxygen"!
Singlet oxygen denotes an excited oxygen molecule and may be written as O2* or O2(a1Δg)!
[Edited on 9-8-2014 by hissingnoise]AJKOER - 9-8-2014 at 23:02
Unionised:
Good point on the wave lengths requirements. So, barring other paths, photolysis at higher elevations (ski resorts, places like Denver,..) may be more
of a possible issue.
However, one does not have to look only to N2O subjected to UV as a possible source of Singlet oxygen itself, or possibly other active oxygen species.
In particular, there may be many possible reactions of oxygen species and radicals, created via photolysis and N2O. One brief affirmative reference,
for example, citing a slow reaction is found at http://pubs.acs.org/doi/abs/10.1021/j100630a005 , which also mentions, that in addition to N2O, O2 itself upon UV exposure can form Singlet
oxygen, as can CO2 and NO2. With respect to radical reactions, see for example http://m.earthobservatory.nasa.gov/Features/ChemistrySunligh... that incidentally cites several pathways to NO2, some involving the formation of a
peroxy radical.
The reason I am suggesting it at all, is also based on my limited experimental observations of the possible increase in flammability of N2O infused
95% Ethanol upon photolysis (see my recent post in SM's Energetic forum). This could be due solely to the richer oxygen source available from N2O in
the vapor phase with the Ethanol, or to the formation of more energetic compounds (like Ethyl nitrite). I did perform a comparison upon ignition of
the UV treated N2O/Ethanol and a more O2 rich Ethanol mix subject to UV, and judgementally gave a slight edge to N2O/Ethanol.
Further photolysis experiments with CH4 and a N2O/O2 mixture may be more tellling on the extent of any photolysis with any observable (by ignition)
removal of the methane.
----------------------------------------------
The reader should note there are currently health warnings over the consumption and misuse of the N2O gas directly (see, for example, http://www.bbc.com/news/uk-28710824 ).
[Edited on 10-8-2014 by AJKOER]unionised - 10-8-2014 at 10:13
"Unionised:
Good point on the wave lengths requirements. So, barring other paths, photolysis at higher elevations (ski resorts, places like Denver,..) may be more
of a possible issue."
No.
Not unless the people of Denver need space suits.
The atmosphere at ground level does contain some reactive oxygen species and NOx
Typical hydroxyl radical concentrations are of the order of a million per cc http://www.atsdr.cdc.gov/toxprofiles/tp65-c5.pdf
Compare that to the bulk gases (about a million million times bigger) and you see that the hydroxyl radicals can't make much NO2
The air in urban areas typically contains abut 10 ppb of NO2.
The reaction with OH radicals could only add about a thousandth of that.
The human body is, and always has been, capable of dealing with these traces of NOx.
So please stop pretending that this is a real issue.
AJKOER - 10-8-2014 at 12:31
Unionised:
OK, I guess phototoxic reactions involving singlet oxygen just cannot exist for the most part, in your opinion.
So I guess this article "Photodegradation and phototoxicity studies of furosemide. Involvement of singlet oxygen in the photoinduced hemolysis and
lipid peroxidation" at http://www.ncbi.nlm.nih.gov/pubmed/9595711 falls in a different category (I have, so far, avoided all that biochemical literature on singlet
oxygen formation). But we are, in fact, still internally consuming the N2O infused beverage, are we not?
But I do agree it is not probably a big problem. Most likely, at worst, just headaches and such, for those with low dose ethyl nitrite exposure or who
knows what. Cheers!
