Sciencemadness Discussion Board

Molten hydroxide as an oxidizing agent.

BromicAcid - 24-5-2004 at 19:07

Molten hydroxide, specifically potassium hydroxide can be quite the oxidizing agent when in contact with air. For example, it will oxidize:

Bi2O3 ----> KBiO3
Fe2O3 ----> K2FeO4
Cr2O3 ----> K2Cr2O7

Might even oxidize chlorate to perchlorate if it did not decompose at those temperatures. Some other high oxidation state compounds probably escape my thoughts at the moment but you get the point. So how about some other reactions involving the use of a molten hydroxide bath which may include molten nitrate added as well. In addition, does anyone know the mechanism of these reactions, they obviously rely on the oxygen in the air but is that the exclusive source of oxidation? I'm just starting to really get into high temperature oxidation, and this isn't the kind of thing they teach you in school any more, this is most definitely old school.

Cyrus - 24-5-2004 at 21:24

Will it do MnO2->K2MnO4?

Mr. Wizard - 24-5-2004 at 22:16

I remember something about fusing MnO2 with sodium carbonate to make Sodium Manganate -MnO3, which will turn to Permanganate upon exposure to oxygen, while in solution. This was often the first step in converting an insoluble unknown to something soluble in quantitative analysis. If I could find my old book I'd verify it.

frogfot - 24-5-2004 at 22:38

There are a synth of KMnO4 from MnO2/KOH/KClO3. The K2MnO4 is then made by boiling KMnO4 in aqueous KOH solution.

I am a fish - 25-5-2004 at 03:06

Quote:
Originally posted by BromicAcid
Might even oxidize chlorate to perchlorate if it did not decompose at those temperatures.


The problem of thermal decomposition could be solved by using a suitable eutectic mixture. A 50-50 mixture of NaOH and KOH melts at 170 °C. This is below the decomposition temperature of KClO3.

[Edited on 25-5-2004 by I am a fish]

BromicAcid - 25-5-2004 at 14:28

Molten hydroxide will convert manganese dioxide to permanganate, forgot that one. Still though I wonder about the reaction, obviously the melt reacts as you end up with a salt of the oxidizer. Perhaps it actually makes the acid anhydride then instantly reacts to the salt under the basic conditions.

In addition how high a temperature do these reactions run at. For example my bismuthate preparation says to fuse KOH with Bi2O3 in in the presence of air. Is just heating to melting and holding it there fused enough? Additional heating could decompose the oxidizing agent even in the molten mix but would it be almost instantly renewed? If so would it act as a catalyst for the decomposition of the hydroxide maybe to explosive decomposition?

[Edited on 5/25/2004 by BromicAcid]

Marvin - 25-5-2004 at 20:07

This is actually more simple than it looks. Reactions which produce pH changes have redox potentials which alter in a big way depending on the pH. In essenance, certain oxidising agents are really powerful, or really whimpy in strong acid or strong base.

Alter the pH enough, and you can do tricks, molten hydroxides are about as extreme a pH change as you can get without resorting to some highly obscure solvents/systems.

In this way permanganate is not possible becuase the ion that uses up the most base is manganate(VI), so on fusing manganese dioxide with a alkali hydroxide atmospheric oxygen (which isnt affected by the pH changes) is a strong enough oxidising agent to produce this. Fortunatly, when you add acid the reverse happens, manganate(VI) becomes strong enough to oxidise itself, and it disproportionates to manganese dioxide and permanganate.

Oxidising chloride, or chlorate to perchlorate doesnt use up base, so I would doubt this redox potential changes enough to be useful.

This is, it in a nutshell, my understanding of it. Ive skipped the [H+], [OH-] redox theory maths, partially because we dont need it to see how altering the equilibrium affects the redox potential, and mainly because when it was explained to me I was busy daydreaming about nitrogen compounds and white phosphorous manufacture.

Theoretic - 26-5-2004 at 02:24

When molten nitrates decompose/are reduced to nitrites, the oxygen atom splits off as... just that, an oxygen atom. In short, it produces atomic oxygen. Could this be the answer to the problem of conversion of chlorate to perchlorate - oxidize with a molten nitrate?
The nitrate ion also produces the oxide ion during reductive decomposition, further accelerating the process...

BromicAcid - 26-5-2004 at 18:53

That was the best explanation yet Marvin, never thought about it like that. In a way it is almost like forcing the equilibrium of an oxidizing agent backward by putting it in the most extreme basic situation and kind of pulling the carpet out from underneath it. Theoretic, your ideas on nascent oxygen are kind of the same as mine for possible oxidation reactions in molten nitrates, along with nitrogen oxides which also oxidize to a degree and possibly the formation of alkali peroxides or superoxides which are extreme oxidizing agents. Anyone know of any books on this subject?

unionised - 27-5-2004 at 14:13

Erm? Hang on a bit? Molten sodium hydroxide is a poor oxidising agent, it will oxidise aluminium, but that is hardly a great achievement.
Oxygen is a very powerful oxidising agent.
A lot of metals are easier to oxidise when they are in the presence of a base.
Think about oxidising Fe++ to Fe+++. You have to remove an electron from something with a positive charge; the electrostatic atraction is working against you. If you convert the stuff to the hydroxide then the molecule, as a whole, has no charge and things get easier.
This is, as Marvin said, related to the fact that [Fe(H2O)6]+++ is a stronger acid than [Fe(H2O)6]++
if you oxidise the iron, you make a stronger acid; if you do that in the presence of a strong base you get a more favourable reaction because you remove the product (an acid) by reaction with the base.
Now, think about a few acids (and, I'm writing the formulae as funny-looking hydrates here to illustrate the point)
Al(OH3), Si(OH)4, P(OH)5, S(OH)6, Cl(OH)7
The higher the oxidation state, the stronger the acid. Higher oxidation states are generally favoured by alkaline conditions. Molten NaOH is about as alkaline as it is easy to get, but it's a piss-poor oxidiser.

[Edited on 27-5-2004 by unionised]

BromicAcid - 30-11-2004 at 10:03

Quote:

Also there is an equilibrium in the melt:

2 OH- <----> H2O + O2-

That helps to contribute to the oxidation action of hydroxides, however this action alone is fairly weak and therefore the actual basicity of the reaction medium is the driving force.


From my thread on sodium bismuthate. At the time I thought this reaction was weak, but I recently was reading upon something that my well prove to be the true oxidizing agent in at least molten potassium hydroxide:
Quote:
Of considerable intrest are the studies which have been reported on the synthesis of KO2 via the reaction:

2KOH + 1.5O2 ---> 2KO2 + H2O

It has been shown that at 210C and 1 atm O2 pressure, that oxidation of KOH yields a product containing 20% KO2 [99]. At 510 C and 1 atm O2 pressure, a product containing 32% KO2 was recovered [84], and after 5 hours at 375 C and 100 atm O2 pressure the product contained 70% KO2 [100].


KO2 is a strong oxidizing agent, and even though oxygen is only roughly 20% of air it is reasonable to assume there is a measureable precentage of KO2 in molten KOH exposed to air due to this reaction. This strong oxidizing agent and the basic enviorment both help to account for the high oxidiation states that can be produced though fusion with KOH.

Quote taken from: Peroxides, Superoxides, and Ozonides of the Alkali and Alkaline Earth Metals. Il'ya Ivanovich Vol'nov

Translated by J. Woroncow Plenum Press 1966

[99] C. Kroger. Z. Anorg. Allgem. Chem. 253:92 (1945)
[84] H. Lux. Z. Anorg. Allgem. Chem. 298:285 (1959)
[100] F. Fisher and H. Ploetze. Z. Anorg. Allgem. Chem. 75:30 (1912)

I don't have acess to any of the references though, they might be interesting reads. It would be somewhat easy to tell if this reaction happens appreciably in a regular air exposure if one had a nickel crucible or other material intert to KOH. KO2 is a yellow solid. Heating clean KOH to 250 C or so and holding it there for an hour and allowing it to cool, the yellow color should at least be noticable.

Edit: Added references and the little trailing bit.

[Edited on 12/1/2004 by BromicAcid]

unionised - 30-11-2004 at 13:53

If that's true then its fascinating but I can't help wondering why I have not heard about it before.

The first of those reactions gives rise to oxide, an even stronger base than a hydroxide melt. That will help a lot of oxidations.

KemiRockarFett - 25-12-2004 at 17:05

Quote:
Originally posted by Marvin
This is actually more simple than it looks. Reactions which produce pH changes have redox potentials which alter in a big way depending on the pH. In essenance, certain oxidising agents are really powerful, or really whimpy in strong acid or strong base.

Alter the pH enough, and you can do tricks, molten hydroxides are about as extreme a pH change as you can get without resorting to some highly obscure solvents/systems.

In this way permanganate is not possible becuase the ion that uses up the most base is manganate(VI), so on fusing manganese dioxide with a alkali hydroxide atmospheric oxygen (which isnt affected by the pH changes) is a strong enough oxidising agent to produce this. Fortunatly, when you add acid the reverse happens, manganate(VI) becomes strong enough to oxidise itself, and it disproportionates to manganese dioxide and permanganate.

Oxidising chloride, or chlorate to perchlorate doesnt use up base, so I would doubt this redox potential changes enough to be useful.

This is, it in a nutshell, my understanding of it. Ive skipped the [H+], [OH-] redox theory maths, partially because we dont need it to see how altering the equilibrium affects the redox potential, and mainly because when it was explained to me I was busy daydreaming about nitrogen compounds and white phosphorous manufacture.


The thing you describe is a simple pourbaix diagram after the researcher Pourbaix who investigated corrosion at different materials.
The diagrams are easily composed by a water area, between 1,23 V and 0 V but due to overvoltage you can add a halv volt. ( Out of the water area the redoxpotential is to high or low for the compound to exist).The diagrams have different areas showing what compound who is dominating under that E0 and pH value.
I can recomend all interested persons to study the diagrams in the book inorganic chemistry written by Gary Wulfsberg. If you search the term pourbaix diagrams in google you may find several of them.

ref 99

S.C. Wack - 14-1-2005 at 11:50

The library is missing 1959.

Attachment: z_anorg_253_92_1945.djvu (116kB)
This file has been downloaded 726 times

unionised - 16-1-2005 at 02:11

Please could you let me know how I can open that file?
Thanks

BromicAcid - 16-1-2005 at 08:14

Here is a viewer for DJVU documents in your web browser. Thank you S.C. Wack for getting that reference, it's fairly interesting.

mick - 16-1-2005 at 08:50

Just caught this post
The acid and alkali idea will work in water, non-aqueous solutions and other systems.
The oxidiser and reducer idea also works in water, non-aqueous solutions and other systems.
Put two elements together and one will be the acid and one will be the base or one will be oxidising and one reducing.
I always liked the idea of the transition temperature. Get the stuff upto just before the transition temperature and you know where you stand.
mick