Orenousername
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Oxidizing properties of N2O
I'm familiar with using hydrogen peroxide to oxidize a stoichiometric solution of alkali iodides and mineral acids to elemental iodine.
Ex. KI + HCl <-> HI + KCl 2HI + H2O2 -> I2 + 2H2O
I'm wondering if nitrous oxide (N2O) would be a strong enough oxidizing agent to perform this same reaction under normal conditions.
Eg. KI + HCl <-> HI + KCl 2HI + N2O -> 2I + N2 +H2O
What do you guys think? Hydrogen peroxide works just fine but I have about about 10kg N2O lying around and if it could be used instead would be
useful..
Also, what about N2O's oxidizing properties in general? I can't find much literature on it outside of organic chemistry. I was able to find some
interesting literature which is probably a bit beyond me though.
http://pubs.rsc.org/en/content/articlehtml/2015/cs/c5cs00339...
[Edited on 5-20-2016 by Orenousername]
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woelen
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N2O is a strong oxidizer, but it is much like oxygen from the air. Only at elevated temperatures it really works as a strong oxidizer, in combustion
processes. At room temperature, the gas is quite inert and does not react easily.
This behavior has to do with its high activation energy. Quite some energy must be put into the molecule before it breaks up and releases that energy
again plus a lot of additional energy in a reaction, which in turn can break up other nearby molecules. Hence, only in fire-like conditions it reacts
easily (many combustible materials also can burn in N2O).
[Edited on 20-5-16 by woelen]
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Orenousername
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Okay, I see. The activation energy could probably be lowered by a catalyst though, no? Decomposition is catalyzed by certain transition metals if
I'm not mistaken. I'm thinking copper or silver for some reason.
I'm thinking that NO2 would probably be a better alternative for most reactions though.
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Zyklon-A
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Loads of catalysts can cause nitrous oxide to decompose, but even if it could be done gently (a big "if"), why would it oxidize iodide? And seriously
how would you setup the experiment? Passing the gas through a catalyst tube first will just send hot oxygen and nitrogen into your pot, at rocket
speeds. I doubt any heterogeneous catalyst will work in water, and if it did, either yields will be bad or they'll be covering your four walls and
still be shit.
No, catalyzing nitrous's decomposition to oxidize something it normally won't is no better than catalyzing peroxide's decomposition for the same
reason. In either case, if the catalyst works, you'll land yourself a hot mess from hell without any desired products. The key is
decomposition is the wrong thing to catalyze.
What you want is something to catalyze iodide's oxidation and I can't help you there. Use peroxide.
BTW most of the catalysts I've heard of for nitrous's decomposition are transition metal oxides.
[Edited on 24-5-2016 by Zyklon-A]
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Orenousername
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Quote: Originally posted by Zyklon-A  | Loads of catalysts can cause nitrous oxide to decompose, but even if it could be done gently (a big "if"), why would it oxidize iodide? And seriously
how would you setup the experiment? Passing the gas through a catalyst tube first will just send hot oxygen and nitrogen into your pot, at rocket
speeds. I doubt any heterogeneous catalyst will work in water, and if it did, either yields will be bad or they'll be covering your four walls and
still be shit.
No, catalyzing nitrous's decomposition to oxidize something it normally won't is no better than catalyzing peroxide's decomposition for the same
reason. In either case, if the catalyst works, you'll land yourself a hot mess from hell without any desired products. The key is
decomposition is the wrong thing to catalyze.
What you want is something to catalyze iodide's oxidation and I can't help you there. Use peroxide.
BTW most of the catalysts I've heard of for nitrous's decomposition are transition metal oxides.
[Edited on 24-5-2016 by Zyklon-A] |
I wasn't suggesting catalyzing decomposition 
Sorry if I was unclear, I was simply speculating that maybe some catalyst will be suitable for direct oxidation. I have read somewhere (can't
remember where) that solutions of iodide are oxidized by molecular oxygen in very small quantities. Maybe shooting hot oxygen and nitrogen into my
pot COULD work 
I understand that it's probably unfeasible, but what am I supposed to do with this 10kg of N2O?
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Zyklon-A
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Gotcha. Yes many iodide salts are oxidized in air but AFAIK group I & II iodides need carbon dioxide to bond to the strong base that would
otherwise make the reaction unfavorable.
Solutions of HI also oxidize forming water and I2, which react further with HI to yield the triiodide ion which darkens the acid.
So it might work to bubble mixed air and carbon dioxide.
Whatever you do with that nitrous, just know its worth way more than the oxygen it contains.
[Edited on 24-5-2016 by Zyklon-A]
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AJKOER
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Reaction of nitrous oxide and water under photolysis (or irradiation) is well known in the art and considered a convenient source of hydroxyl
radicals:
N2O + H2O + hv → N2 + •OH + OH-
Source: See, for example, https://books.google.com/books?id=Om_TKidEjToC&pg=PA229&...
Relatedly, the action of O2 with NaOH at around 100 C with a touch of a Mg salt (the created Mg(OH)2 providing a surface interface catalyst) can also
form hydroxyl radicals and related oxygen species which serve as the commercial basis for the bleaching of wood plup. The process is referred to as
oxygen-alkali bleaching and I suspect is based on a reversal of the known radical reaction below using higher temperature and pressure in presence of
hydroxyl radical scavengers (like ferrous and cuprous,..) and possibly a catalytic organic or metal based intermediate:
O2•- + •OH = O2 + OH- (Reaction 49 in Table 2.1, page 23 at https://www.google.com/url?sa=t&source=web&rct=j&... )
Also, created hydroxyl radical can react, for example, with friendly ethyl alcohol in the presence of dioxygen (O2), forming the superoxide anion
(.O2-) as follows:
CH3CH2OH + •OH → H2O + •CH3CHOH
•CH3CHOH + O2 → CH3CHOH + H+ + O2•-
Source: See, for example, full text available at: https://www.researchgate.net/publication/222191508_Kinetics_...
In general, the hydroxyl radical can breakdown large organic compounds (like dyes and stains) into more elementary compounds (like CO2 and H2O).
The O2•- species can further react with say, hydogen peroxide, in the so called Haber-Weiss (HW) reaction, which is fast in the presence of copper
and iron ions (which consume the hydroxyl radicals and/or the hydroxyl ions):
O2•- + H2O2 → •OH + OH- + O2 (see https://www.google.com/url?sa=t&source=web&rct=j&... )
My claim of the reversal of the radical reaction at higher temperature and pressure for plup bleaching appears to by supported by a study (see https://www.google.com/url?sa=t&source=web&rct=j&... ) that pre-impregnated wood plup with H2O2 (as opposed to post) and then applied
oxygen-alkali bleaching, thereby obtaining superior results. As my postulated products are •OH and O2•-, the latter can further react via HW
(assuming a transition metal presence in the plup) to produce more •OH and recycle the O2 input.
The Haber-Weiss reaction is probably, in my opinion, the most cited near non-reaction, as the need for an appropriate transition metal catalyst is
not generally detailed.
Also, other aggressive oxidizing species are possibly created with oxygen in the persistent presence of hydroxyl radicals at an appropriate pH
including, for example:
•OH + O2 ⇄ HO3• pKa -2.1 (see http://onlinelibrary.wiley.com/doi/10.1002/poc.1812/abstract )
And, of most curiosity to me, is the cited interaction between nitrous oxide and the superoxide radical anion in the presence of a catalyst:
N2O + O2•- → N2 + O3•-
Reference: See page 151at https://www.slideshare.net/mobile/drboon/catalytic-decomposi... and also page 879 of "Handbook of Advanced Methods and Processes in Oxidation
Catalysis: From ...", edited by Daniel Duprez, Fabrizio Cavani, available at https://books.google.com/books?id=-tG3CgAAQBAJ&pg=PR5&am... by selecting page 877 from page menu and proceeding manually to p. 879.
which, would make one strong oxidizing mix. [Edit] At least, one would so think, given its relation to ozone. However, per the Duprez's Handbook, on
page 258, the ozonide radical is described as short lived and relatively inert intermediate decomposing into OH- and O2.
[Edited on 26-5-2016 by AJKOER]
[Edited on 26-5-2016 by AJKOER]
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Orenousername
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Wow, great information here AJKOER. Wasn't exactly planning on dabbling in free radical chemistry yet, but it looks like I don't have much of a choice
in the matter anymore
The superoxide anions are of particular interest to me, do you think that it would be plausible to prepare KO2 from N2O in solution with water and
ethanol in the presence of K+? Wikipedia says that KO2 decomposes in water into H2O2 and the corresponding hydroxide.. Is there a practical way this
can be overcome in the amateur lab? Hydroxides and H2O2 are still useful however, but having some superoxide salts would be pretty cool.
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AJKOER
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Orenousename:
Research preparations of MgO2, CaO2,...and one will observe the use of H2O2. The latter in the presence of OH- and light, forms HO2- which is
apparently even more light sensitive than H2O2 which also forms hydroxyl radicals in strong light. The interaction of the hydroxyl radical and H2O2
creates the HO2 radical.
So, what I am suggesting, is that akaline H2O2 and light, is like working with HO2- and HO2 radical.
[Edited on 28-5-2016 by AJKOER]
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unionised
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| Quote: Originally posted by AJKOER | Reaction of nitrous oxide and water under photolysis (or irradiation) is well known in the art and considered a convenient source of hydroxyl
radicals:
N2O + H2O + hv → N2 + •OH + OH-
|
That reaction is clearly wrong because the charges don't balance.
You have messed up the copying and forgotten the solvated electron.
Since creation of those electrons takes rather "forcing" conductions, it's hard to see this as "convenient" (unless you happen to have a 10 MeV
linear accelerator of course)
Do you have any references for your assertion that
"Relatedly, the action of O2 with NaOH at around 100 C with a touch of a Mg salt (the created Mg(OH)2 providing a surface interface catalyst) can also
form hydroxyl radicals and related oxygen species"?
[Edited on 28-5-16 by unionised]
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AJKOER
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Unionised:
Thanks, I agree a little sloppy, but the source of the solvated electron, in this case, is light.
Other sources are microwave radiation or the employment of a radiative element.
Rewriting as:
N2O + e-(aq) ---hv--) N2 + .OH + OH-
is also possible but some will ask where is the hydrogen and added oxygen coming from? Answer (aq) means water.
------------------------------------------------
Here is a reference that outlines the underlying science of plup bleaching which I view as more of an art than pure science. Link to pdf:
https://www.google.com/url?sa=t&source=web&rct=j&url=http://www.sefs.washington.edu/classes.pse.476/Readings/Chemical%2520Pulping%2520A137
_A146.pdf&ved=0ahUKEwibjOOZ-YPMAhVF9h4KHUpEC74QFggvMAg&usg=AFQjCNGBqnuEtk6JKDevNkwyTo4qgeY0dQ&sig2=Kz38ElcLD6gjMIjU0MpScw
[Edited on 28-5-2016 by AJKOER]
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deltaH
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AJKOER, so basically a single electron reduction of N2O in a protic solvent transforms it into an oxidant only slightly weaker than
fluorine?! 
The reduction potential of .OH is, after all, a massive 2.8V! (http://www.h2o2.com/products-and-services/us-peroxide-techno...)
If so, there are more than one way to skin this cat.
Perhaps one could take a leaf out of Fenton's book and use an acidified solution of a ferrous salt to act a single electron reductant?
Thus
e.g. Fe2+(aq) + N2O(g) + H+(aq) => Fe3+(aq) + N2(g) + .OH(aq)
Would be an easy way for OP to carry out wet oxidations with his N2O in an amateur setting and OTC reagents.
Orenousername, if you'd like some background to where my idea comes from, it's adapted from Fenton Chemistry which is used to
generated hydroxyl radicals from hydrogen peroxide using acidified ferrous salt solutions. The difference here though is that you wouldn't be using it
in catalytic amounts, but stoichiometrically.
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Orenousername
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| Quote: Originally posted by AJKOER | Orenousename:
Research preparations of MgO2, CaO2,...and one will observe the use of H2O2. The latter in the presence of OH- and light, forms HO2- which is
apparently even more light sensitive than H2O2 which also forms hydroxyl radicals in strong light. The interaction of the hydroxyl radical and H2O2
creates the HO2 radical.
So, what I am suggesting, is that akaline H2O2 and light, is like working with HO2- and HO2 radical.
[Edited on 28-5-2016 by AJKOER] |
MgO2 and CaO2 are peroxide salts, no? Does O2- react with water to form HO2- and OH? This does seem very likely to my (uneducated) mind. In which
case, only peroxide salts could be formed from the hydroxyl radical in aqueous solution, right? Although, if I'm not mistaken peroxide salts
decompose rapidly in aqueous solution as well..
[Edited on 5-29-2016 by Orenousername]
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AJKOER
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| Quote: Originally posted by deltaH | AJKOER, so basically a single electron reduction of N2O in a protic solvent transforms it into an oxidant only slightly weaker than
fluorine?!
The reduction potential of .OH is, after all, a massive 2.8V! (http://www.h2o2.com/products-and-services/us-peroxide-techno...)
If so, there are more than one way to skin this cat.
Perhaps one could take a leaf out of Fenton's book and use an acidified solution of a ferrous salt to act a single electron reductant?
Thus
e.g. Fe2+(aq) + N2O(g) + H+(aq) => Fe3+(aq) + N2(g) + .OH(aq)
Would be an easy way for OP to carry out wet oxidations with his N2O in an amateur setting and OTC reagents.
.....................
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Technically, in my research, the correct and qualified reaction appears to be observed in heterogeneous conditions:
Fe(ll)L + N2O (aq) + H+(aq) --Catalyst-> Fe(lll)L + N2 (g) + .OH (aq)
where the reaction is likely selective both to the availability of the appropriate ligand 'L' and possible transition metal catalyst.
As to an example of a heterogeneous environment, in the human body, to quote:
"We show that N2O indirectly inhibits T-type calcium channels through free radical reactions. These reactions depend on a histidine residue on the
channel that binds metal ions."
Also:
"It is important to note that the rates of adrenochrome production in H2O2 + FeCl3 and N2O + FeCl3 are very similar for the first 10 min. "
where adrenochrome assay is a standard for reactive oxygen species (ROS) production (per Valerino & McCormack, 1971).
Link: https://physoc.onlinelibrary.wiley.com/doi/pdf/10.1113/jphys...
A better choice of a transition metal than Fe would be Ti (or Sc , V) with respect to a charge transfer reaction (see https://pubs.acs.org/doi/abs/10.1021/jp004613%2B ).
[Edited on 27-5-2018 by AJKOER]
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unionised
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Any evidence that this reaction actually happens?
Fe(ll)L + N2O (aq) + H+(aq) --Catalyst-> Fe(lll)L + N2 (g) + .OH (aq)
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Tsjerk
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| Quote: Originally posted by unionised | Any evidence that this reaction actually happens?
Fe(ll)L + N2O (aq) + H+(aq) --Catalyst-> Fe(lll)L + N2 (g) + .OH (aq)
|
I'm pretty sure there is none, because it will not happen. I would be more than happy if it would happen as I would be able to publish it in Science
or Nature.I know AJOEKER didn't pull this from thin air but It is, as before, pulled completely out of context.
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AJKOER
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The section title, on page 142 of the article, 'Free radical signalling underlies inhibition of CaV3.2 T-type calcium channels by nitrous oxide in the
pain pathway', relating to the adrenochrome test noted above for reactive oxygen species (which could include hydroxyl radicals) is:
'N2O causes the generation of ROS in the presence of transition metal ions'.
There is also a soil study, 'Iron oxidation affects nitrous oxide emissions via donating electrons to denitrification in paddy soils', available at http://wwwuser.gwdg.de/~kuzyakov/Geoderma_2016_Wang-Milan_Fe... , to quote from the abstract:
"The soil with high Fe(II) concentration emitted less N2O than did the other soil with low Fe(II) concentration. "
Note, both studies are performed in very heterogeneous conditions, and I suspect repeating in a homogeneous environment may prove to be disappointing
without the proper metal catalyst (like possibly Ti, V, traces of Co,...), surface chemistry or ligands or bacteria or even perhaps, in my
speculation, phosphates promoting solvated electron formation (note the title includes the words 'donating electrons').
[Edited on 28-5-2018 by AJKOER]
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AJKOER
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I apparently overlooked the most obvious catalyst in the consumption of N2O in wet iron rich soils from room light to sunlight!
See 'Evidence of singlet oxygen and hydroxyl radical formation in aqueous goethite suspension using spin-trapping electron paramagnetic resonance
(EPR)' by Sang Kuk Han, et al, link: https://www.academia.edu/15600734/Evidence_of_singlet_oxygen...
In the current context, goethite can be formed from added FeSO4 by simultaneous oxidation (from say a slow electrochemical oxidation in the presence
of O2, CO2 and an electrolyte, or faster from any in situ created H2O2 in the presence of transition metals, organic acid, dissolved oxygen and light)
and precipitation with carbonate ion (source: see, for example, https://onlinelibrary.wiley.com/doi/pdf/10.1002/jctb.5020250... ). Then, with light energy, some solvated electron formation:
Fe(lll)O(OH) + hv → photogeneration of electron-hole (e-, h+) → ROS (•O2-, •OH)
followed by:
N2O + e- (aqueous/surface) --> N2 (g) + •O- (aq)
•O- + H2O = •OH + OH-
So, one could also possibly view this as a photolysis assisted degradation of N2O in visible light upon adding a ferrous salt.
The authors of the cited paper also largely ignored mention of any possible causality associated with sunlight, with the exception of one reference:
'Melton, E.D., Schmidt, C., Kappler, A., 2012. Microbial iron (II) oxidation in littoral freshwater lake sediment: the potential for competition
between phototrophic vs.nitrate-reducing iron (II)-oxidizers. Front. Microbiol. 3.'
referring to 'phototrophic' microbes.
I am, to be fair, more at fault having recently noted:
"And the presence of any photoactive metal oxides (like ZnO, MgO, TiO2,..) releasing electrons in sunlight may replace the need for UV light.
Reaction:
e- (aq) + N2O --> N2 + •O-
•O- + H2O = •OH + OH- "
Link: http://www.sciencemadness.org/talk/viewthread.php?tid=82541#... and a link to the literature employing, for example, ZnO/N2O: https://pubs.acs.org/doi/abs/10.1021/j100657a003 .
-------------------------------------------------------------
Just noticed that the reported soil samples were incubated in the dark for 40 days at 25 C.
As such, my claim of any possible photolysis effect is, for this study, withdrawn, as the argument that there could be residual photocatalytic
produced products in the soil is weak as the soil samples were also air-dried. The first cited animal study involving N2O was conducted apparently
under laboratory lighting, but for short durations (under an hour).
Note, performing the study in the dark could be viewed as advisable to limit any possible photocatalytic effect.
[Edited on 30-5-2018 by AJKOER]
[Edited on 30-5-2018 by AJKOER]
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