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Potassium oxide decomposition
Potassium oxide decomposes at 300 °C, but what are decomposition products?
By the way, I found an easy way to make it: Decomposition of K2S2O5 at 190°C to K2O and SO2. K2S2O5 is available as wine additive.
[Edited on 29-6-2010 by Random]
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woelen
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No, you are not correct. What you write about decomposition of K2S2O5 sounds too good to be true. I can imagine that K2S2O5 decomposes giving SO2 and
K2SO3, but the second molecule of SO2 definitely is not given off at 190 C. That probably would require 1000+ C. If this were possible then we would
have a fantastic source of a strong base like K2O, but that is not the case.
Potassium oxide does not decompose at 300 C, but it can absorb oxygen, giving rise to formation of K2O2 or even KO2.
[Edited on 29-6-10 by woelen]
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DJF90
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You can work out temperatures of decomposition using free energy cycles.
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Random
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Quote: Originally posted by woelen  | No, you are not correct. What you write about decomposition of K2S2O5 sounds too good to be true. I can imagine that K2S2O5 decomposes giving SO2 and
K2SO3, but the second molecule of SO2 definitely is not given off at 190 C. That probably would require 1000+ C. If this were possible then we would
have a fantastic source of a strong base like K2O, but that is not the case.
Potassium oxide does not decompose at 300 C, but it can absorb oxygen, giving rise to formation of K2O2 or even KO2.
[Edited on 29-6-10 by woelen] |
I am not sure about decomposition of potassium metabisulfite, but this is what I read on wikipedia:
| Quote: |
Potassium metabisulfite has a monoclinic crystal structure which decomposes at 190°C, yielding potassium oxide and sulfur dioxide:
K2S2O5(s) → K2O(s) + 2SO2(g)
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blogfast25
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I can't see this being correct either. Woelen's most pragmatic statement 'if this were possible then we would have a fantastic source of a strong base
like K2O, but that is not the case' is the real clincher here: if K2O could be made easily from an OTC, I'd be making potassium via thermite or by
reduction by Mg w/o the nasty hydrogen flying around from KOH...
And Na, from Na2O from Na2S2O5 (the more readily available metabisulphite)... I've got some of that and might run a test tube test... Mind the SO2!
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Random
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| Quote: Originally posted by blogfast25 | I can't see this being correct either. Woelen's most pragmatic statement 'if this were possible then we would have a fantastic source of a strong base
like K2O, but that is not the case' is the real clincher here: if K2O could be made easily from an OTC, I'd be making potassium via thermite or by
reduction by Mg w/o the nasty hydrogen flying around from KOH...
And Na, from Na2O from Na2S2O5 (the more readily available metabisulphite)... I've got some of that and might run a test tube test... Mind the SO2!
|
I am really interested in results of the experiment, please post them if you'll do it 
By the way, decomposition products may not be same like with potassium metabisulfite decomposition.
[Edited on 29-6-2010 by Random]
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JohnWW
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Actually, "potassium oxide, K2O", even if it could somehow be obtained as a "technical grade" reagent, is liable to contain also peroxide and
superoxide, K2O2 and KO2, particularly by direct combination of the elements, and the ozonide KO3 is also known. On exposure to ordinary air at
ordinary temperatures, it would immediately start absorbing H2O vapor and CO2 to form KOH and KHCO3 and K2CO3, along with some KOOH.
Besides, you CANNOT simply heat K2O in a glass vessel to any elevated temperature - it would simply dissolve the glass to form largely K2SiO3. The
only really safe common materials for a vessel to heat the stuff in would be either pure Ni, or one of the high-nickel Inconel or Monel alloys.
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Anders Hoveland
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Anhydrous Na2SO3 does not melt until 500°C. Melting point, not the decomposition.
Now the nitrites might work better.
Potassium nitrate does not decompose until 440°C, so I highly doubt that wikipedia was correct about a decomposition at a much lower temperature for
the bisulfite. Actually NaNO2 has a 271 °C decomp, accourding to wikipedia. If this is true, than KNO2 would probably be similar and 271C is pretty a
reasonable temperature to make K2O.
2KNO2 --> K2O + NO2 + NO
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Polverone
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All these materials may decompose in that they become mixtures instead of pure substances, but I very much doubt they're going to leave oxides.
Also, use some critical judgment about sources: an un-cited, astonishing bit of chemistry found on Wikipedia is almost certainly an error rather than
a breakthrough. Lange's Handbook of Chemistry and the CRC Handbook of Chemistry and Physics both contain melting/decomposition points for thousands of
substances and can be found inexpensively second-hand, or even as digital copies if you're penniless.
At red/orange heat K and Na nitrates and nitrites do not decompose to the oxide, at least not at any appreciable rate. Heated strongly in the presence
of air and absence of additional oxidizing or reducing material, the nitrites and nitrates both become a nitrate/nitrite mixture (more nitrate than
nitrite). After cooling the residue is only mildly alkaline in water solution. This is from personal experience; I too once hoped that I might be able
to make very strong bases with little more than fertilizer and a furnace.
PGP Key and corresponding e-mail address
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Formatik
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Some reactions known to give Na2O: glowing NaNO3 with MnO2 under absence of air (Lieb. Ann. 119 [1861] 375). Glowing here should be in the range
around 400 to 600 C. That's most likely the lowest you will get without something like a starting material of alkali metal or peroxides. Molten NaCl
solubilizes BaO forming a melt which strongly attacks Pt. It's probable, the following reaction occurs: BaO + 2 NaCl = BaCl2 + Na2O (Z. phys. Ch. 78
[1912] 567). Reaction of gaseous NaCl and CO, forms Na2O.
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woelen
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| Quote: Originally posted by Anders Hoveland | Potassium nitrate does not decompose until 440°C, so I highly doubt that wikipedia was correct about a decomposition at a much lower temperature for
the bisulfite. Actually NaNO2 has a 271 °C decomp, accourding to wikipedia. If this is true, than KNO2 would probably be similar and 271C is pretty a
reasonable temperature to make K2O.
2KNO2 --> K2O + NO2 + NO |
This is not true, I have heated NaNO2 in a supremax test tube and it simply melts, even when very hot. It does not decompose. I once tried this in een
effort to obtain dry NO and NO2 without any water vapor, but that does not work. I'm not sure about KNO2 but probably this is the same.
Alkali nitrites are very stable with respect to heat and are not easily decomposed.
Maybe it is a good thing to fix some errors on Wikipedia? This kind of wrong information is not good at all. Sometimes, when I read info about
chemicals on Wikipedia and I find an error, I change it.
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We should fix those errors on wikipedia, if they are not true. Though, something is happening at range of 150-190°C with this compound. Just we don't
know which decomposition products are. Maybe some so2 could react with K2O formed? If K2SO3 is formed, we can react it with water to form K+ ions, H2O
and SO2. After that we could add Ca(OH)2 to make KOH and Ca ions.
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woelen
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If there is decomposition of K2S2O5 at 190 C then the only reasonable thing is the following reaction:
K2S2O5 --> K2SO3 + SO2
Nothing else can happen. So, if there is decomposition, then we can be quite sure about the reaction.
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Anders Hoveland
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There is always CaO (lime) that is available at HomeDepot.
Perhaps using Barium oxide in a molten mixture?
BaO + K2SO4 --> BaSO4 + K2O
Barium sulfate has a signif. higher mp so the molten K2O might separate out as a liquid with only minimal BaSO4 dissolved in it.
CH3C(NH)NHCH2CH3 + K2O --> KNC(NH2)CH3 + KOCH2CH3
If one of you did succeed at making K2O, then the above reaction might be a way to make the useful ethoxide, for example if you were making DADNE.
CH3C(NH)NHCH2CH3 probably could be made from anhydrous NH3 and ethyl acetate.
Of course there is always 2KN3 + 2KNO2 --> 2K2O + 2NO + 3N2
if you are really desperate, but it is so difficult to prepare the azide, that you might as well make metallic K electrolytically.
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Formatik
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I looked recently at about one and a half pages of decomposition of K2S2O5 in Gmelin. No K2O forming as a product. Long story short: only, K2S2O3,
K2SO3, K2SO4, SO2, S, and K-polysulfide have been characterized as thermal decomposition products. Proving what most of us already had a hunch for,
the wikipedia article is wrong. Also not surprising, there's no reference that was given for the supposed decomposition.
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franklyn
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Decomposition in this case is a very poor descriptor as it implies
something is given off , which cannot be the case. A possible
outcome is the combination with atmospheric oxygen to form the
peroxide. The oxides of Potassium are , K2O , K2O2 , KO2 , KO3 ,
and decomposition when it occurs is toward formation of K2O.
At the indicated temperature the solid K2O may sublimate,
the likely reason no value is given for boiling point.
http://www.webelements.com/compounds/potassium/dipotassium_o...
http://www.webelements.com/compounds/potassium/dipotassium_p...
http://www.webelements.com/compounds/potassium/potassium_sup...
http://www.chemguide.co.uk/inorganic/group1/reacto2.html
http://www.springerlink.com/content/w30707824j61m061
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Formatik
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| Quote: Originally posted by franklyn | Decomposition in this case is a very poor descriptor as it implies
something is given off , which cannot be the case. A possible
outcome is the combination with atmospheric oxygen to form the
peroxide. The oxides of Potassium are , K2O , K2O2 , KO2 , KO3 ,
and decomposition when it occurs is toward formation of K2O.
At the indicated temperature the solid K2O may sublimate,
the likely reason no value is given for boiling point. |
Even if the compound is being decomposed in a good current of air and combined with it, it will still rather form a compound like K2SO4 than K2O.
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smuv
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Literature...ugh...
I am resurecting this thread because I found the article that I suspect wikipedia failed to reference. In this article they DO say that K2O is fomed
around 190c, but the data seems fishy. First the TG trace graph's Y axis is labeled wrong, with 100c showing 100% mass lost, and 300c showing about
77% mass lost. I assume they simply mean mass not mass lost. But then 100-76--> 24% loss, and K2O to K2S2O5 requires a much higher loss. Making
things even more confusing is, forming K2SO3 would be a 29% mass loss. So if I am interpreting their data right, that didn't even happen. Their data
actually works out better if you conclude that they let in air (instead of the "99.9995% nitrogen atmosphere reported) and potassium
sulfate was formed than any of the other possibilities.
Bullshit bullshit bullshit.
I got excited when I first read this article, because they said in the intro unequivocally that the decomp was one step: K2S2O5 --> K2O + 2SO2.
But after reviewing their data, it seems like shoddy work and doesn't even make sense. Attached are two papers, by the same group, reporting the same
shoddy data...in different journals. What bullshit;
B. Jankovic, you are a terrible chemist, and you just wasted my time. 
Attachment: bullshit1.pdf (805kB)
This file has been downloaded 4100 times
Attachment: bullshit2.pdf (236kB)
This file has been downloaded 4301 times
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BromicAcid
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Ahh, this is actually a good one 
From Peroxides, Superoxides, and Ozonides of Alkali and Alkaline Earth Metals:
| Quote: | ...for example, [potassium peroxide] can be produced by heating potassium oxide, K2O, in a vacuum (10^-5 mm Hg) at temperatures above 450°C, i.e.,
2K2O (Solid) -----> K2O2 (Solid) + 2K (gas) |
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IndependentBoffin
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After several years of absorbing content from technical papers as my hobby/job, even now when I read papers full of word spaghetti, unless the paper
comes highly recommended, it seems that the author(s) just can't be bothered to understand their content and speak plainly (in which case why should I
waste my time).
After reading a paper you always think to yourself what you understood and took away from reading it.
[Edited on 23-6-2011 by IndependentBoffin]
I can sell the following:
1) Various high purity non-ferrous metals - Ni, Co, Ta, Zr, Mo, Ti, Nb.
2) Alkex para-aramid Korean Kevlar analogue fabric (about 50% Du Pont's prices)
3) NdFeB magnets
4) High purity technical ceramics
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KonkreteRocketry
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interesting, but I have K2O as a product in a rocket chamber at a temperature of around 1900 degree, what happens to the K2O ?
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Endimion17
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Nothing.
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Ral123
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Can some peroxide containment be achieved by reacting alkali carbonates/nitrites/hydroxide with magnesium, to produce the oxide and some excess metal,
then the product be heated in oxygen. After all that, will some H2O2 be formed in HCl?
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woelen
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The K2O will not exist for long. If it is formed, then it will be accompanied with CO2, SO2 and who knows what more combustion products (depending on
what is used as oxidizer and what fuel you use). It soon will combine with these to K2SO3, K2CO3 and otherwise it will combine with water and CO2 from
the air. K2O cannot exist in nature, it is too reactive and only can exist in a tightly sealed bottle under perfectly dry air, free of CO2.
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KonkreteRocketry
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| Quote: Originally posted by woelen | | The K2O will not exist for long. If it is formed, then it will be accompanied with CO2, SO2 and who knows what more combustion products (depending on
what is used as oxidizer and what fuel you use). It soon will combine with these to K2SO3, K2CO3 and otherwise it will combine with water and CO2 from
the air. K2O cannot exist in nature, it is too reactive and only can exist in a tightly sealed bottle under perfectly dry air, free of CO2.
|
hehe yes, but here it is only nitrogen gas, magnesium oxides out there, no CO2 or SO2 or any O2 or anything else, so the K2O is alone, what will
happen then ? become a gas ?
Also, what will happen to K2CO3 at 1900 degree ? I see that it decomposes to K2O and CO2 ... ummm ? so if at 1900 degree, K2O will be alone ?
[Edited on 5-5-2013 by KonkreteRocketry]
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