John paul III - 14-6-2018 at 15:16
I came across an interesting reaction here: http://www.allreactions.com/index.php/group-1a/potassium-k/p...
KOH + (2-4) H2O2 - > K2O2 * (2-4) H2O2 at 0°C
Obviously that doesn't seem right since the reaction is unbalanced (I guess the product should be a hydrate).
Is this reaction real? And why does it have to
be at 0°C? (does it?)
DoctorPhilosophy - 14-6-2018 at 17:19
I see no practical difference, whether it is peroxyhydrate, K peroxide or H peroxide, in any case the active ion is peroxide or oxygen in 1- oxidation
state. Or simply said oxidant. Theoretically tough, I am curious, how can we know? Does such test exist? Similar example is peroxide + sulfate, or
persulfate? How are such things verified to be something we think it is?
John paul III - 15-6-2018 at 04:15
One only needs to heat up the product to 500°C (hopefully the cocrystal H2O2 evaporates/decomposes slowly enough, if it's H2O2 not water). At this
themperature k2o2 decomposes to k2o and at 340 - 430 °C K2O disproportionates to K and K2O2. The potassium metal should become visible (or catch fire
with air) and at lower temp
K2O2 and K probably recombine explosively. Still, testable in small quantity. If someone has conc. H2O2 at hand and is willing to try (think about
potassium metal!) I'll be very grateful
[Edited on 15-6-2018 by John paul III]
DoctorPhilosophy - 15-6-2018 at 05:24
First time I hear that K2O disproportionates. Although I heard that for Fe oxides. Hope somebody has conditions neccessary to try that. My hands are
kind of tied right now, with invisible handcuffs or soemthing.