AJKOER - 26-8-2022 at 17:52
Per this educational reference (link: https://www.youtube.com/watch?v=y-uSwAisad8 ), a hypothetical (?) reaction of Laughing Gas (N2O) on hypochlorite (OCl-) proceeds as follows:
Basic conditions:
N2O + 2 OCl- + 2 OH- -?-> 2 NO2- + H2O + 2 Cl-
Acidic conditions:
N2O + 2 OCl- + H2O -?-> 2 NO2- + 2 H+ + 2 Cl-
which implies with HOCl, one has:
N2O + 2 HOCl + H2O -?-> 2 HNO2 + 2 HCl
And as nitrous acid is unstable, one might thus expect a release of NO and NO2 gases, and with added oxygen presence (per limited addition of H2O2 to
the HOCl pre-addition of the Laughing gas), essentially NO2 as sourced (?!) from safe N2O Laughing Gas.
More crazy, I actually employ N2O in a NaOCl/HOCl system once, actually as a path to chlorate under strong sunlight (see http://www.sciencemadness.org/talk/viewthread.php?tid=34429#...), where I do recall the system did appear to be somewhat more light sensitive than
I expected (note, nitrite is a more potent photocatalyst than nitrate).
I also know (see Table 2.1 here http://www.diva-portal.se/smash/get/diva2:730237/FULLTEXT01....) that:
N2O + e-(aq) -> N2 + .O-
.O- + H2O = .OH + OH-
.H = H+ + e-
.H + N2O -> N2 + .OH
H+ + .O- = .OH
with the N2O + e- reaction moving to the right except at very high pH, but in the cited opening thread postulated reaction above, there is
interestingly, no claimed creation of nitrogen gas.
Any support for the claimed reaction with Laughing Gas acting on hypochlorite?
[Edited on 27-8-2022 by AJKOER]
AJKOER - 26-8-2022 at 18:53
Actually, I think I just have an answer for N2O dissolved in transition metal rich water with dissolved oxygen, starting with a transition metal M in
the presence of a hypochlorite, commencing with a metal auto-oxidation reaction:
M -> M+ + e-
M+ -> M++ + e-
O2 + e- = .O2-
N2O + e- -> N2 + .O-
.O- + H2O = .OH + OH-
.OH + ClO- -> OH- + .OCl
.O2- + .OCl -> O2 + OCl-
My expected Net Reaction:
M + N2O + H2O -- OCl-, O2 --> M++ + N2 + 2 OH-
So, in the presence of a transition metal, oxygen and hypochlorite, I would expect the N2O (aq) to be possibly subject to consumption liberating just
elemental nitrogen gas and elevating the pH.
Here are supporting comments from the ACS (https://pubs.acs.org/doi/full/10.1021/jacs.9b08241) to quote:
"Despite its utility as an oxygen-atom transfer reagent for transition metals, nitrous oxide (N2O) is a notoriously poor ligand, and its coordination
chemistry has been limited to a few terminal, end-on κ1-N complexes. Here, the synthesis of a mononuclear cobalt complex possessing a side-on-bound
N2O molecule is reported."
Another reference: "Mechanistic Studies of Oxygen-Atom Transfer (OAT) in the Homogeneous Conversion of N2O by Ru Pincer Complexes" here https://www.mdpi.com/2304-6740/10/6/69/htm .
Note: Reading the references make clear these aqueous reaction paths requires very particular transition metal complexes. However, photolysis of a gas
mix of, for example, H2 and N2O, with the right light, will readily explode. In general, N2O is viewed as problematic stable greenhouse gas, albeit,
capable of depleting ozone.
[Edited on 27-8-2022 by AJKOER]
Keras - 26-8-2022 at 23:18
The solubility of N₂O in water is not that great. Even if that reaction is theoretically possible, you might need tons of gas to synthesis the least
amount of sodium nitrite…
N₂O is much more soluble in fat, and that’s why it is used as a propellant gas in whipped cream canisters.
AJKOER - 27-8-2022 at 07:24
interestingly, cold water solutions at 5 C, N2O is depicted as highly soluble (see https://cdnsciencepub.com/doi/10.4141/cjss2012-021 ) and also has been described a non-ideal gas. To quote:
"Nitrous oxide is highly soluble in water (at 5°C 1.0 mL N2O-N mL−1 water=0.0425 mols N2O L−1 water) (Dowdell et al. 1979; Davidson and Swank
1990; Heincke 2001), with its solubility increasing as temperature decreases (Weiss and Price 1980; Heincke and Kaupenjohann 1999). During the winter
period in northern latitudes, there is the potential for the formation of continuous ice layers that can restrict the diffusion of N2O to the
atmosphere, trapping N2O at a depth where it can dissolve in cold water (Davidson and Swank 1990; Burton and Beauchamp 1994)."
Keras - 27-8-2022 at 07:51
Wikipedia says 1.5 g/L @ 15 °C. That’s 34 mmol, or 0.76 L (considering 1 mol = 22.4 L).
Not bad, but 34 mmol/L, good luck for obtaining any substantial quantity of nitrite (34 mmol of sodium nitrite represents 2.3 g, and that if the yield
is 100%…)
[Edited on 27-8-2022 by Keras]
AJKOER - 27-8-2022 at 08:19
Did uncover a possible consistent photolysis path, per a source (see http://www.ccpo.odu.edu/SEES/ozone/class/Chap_5/5_4.htm ), the following reactions:
N2O + hc/lambda --> N2 + .O(1D)
N2O + .O(1D) --> 2 .NO
And as the action of UV on HOCl proceeds as follows:
HOCl + hc/lambda --> .HO + .Cl
followed by:
.HO + .NO -> HONO
the associated creation of nitrous acid, which is in accord to quote form the opening thread:
"which implies with HOCl, one has:
N2O + 2 HOCl + H2O -?-> 2 HNO2 + 2 HCl "
albeit, the conditions are quite different (namely, as gases, or gas/liquid interaction, under strong UV exposure) from my initial expectation.
However, in my cited chlorate preparation thread, I did employ an aluminum filament in a microwave to induce a high temperature plasma arc converting
resident air to NO also, followed by a low level photolysis with HOCl, so the chemistry may actually be possible for limited nitrite creation.
[Edited on 27-8-2022 by AJKOER]
Keras - 27-8-2022 at 09:34
Well, I must apologise. I mean what I said was border idiotic. Normally, the N₂O is consumed by the reaction, which turns it into NO₂⁻ which is
very soluble, so you can dissolve more as the nitrite is formed. Try bubbling nitrous oxide very slowly into cold bleach and see what you get. It
might be difficult to separate nitrite from residual hypochlorite, though.