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Author: Subject: Strenght of oxidizers
woelen
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[*] posted on 29-11-2005 at 01:04


Quote:
Originally posted by unionised
Phone NASA and tell them to cancel the next shuttle- perchlorates aren't fast oxidisers any more.

Please read this thread more carefully :mad:. We are talking here about redox reactions in solution, in the cold, not about reactions of solids at high temperatures. So your comparison is not a valid one.
Perchlorates in the cold, dissolved in water, simply are NOT good oxidizing agents, evan at concentrations as high as 60%. Try it yourself!

And the list I compiled can be useful. E.g. a compound like Na2S2O8 can be very useful for analysis purposes, but for synthesis purposes it is less useful (although not useless), simply because its reactions are so sluggish.

The list indeed must not be regarded as a scientificly founded list as the electrochemical series, it is more an indicative list. With some chemical reactions, the positions of individual chemicals can be swapped, but the list gives a fairly good overall impression. That is all what I claim, not more, not less.




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unionised
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[*] posted on 29-11-2005 at 11:27


I read this far
"I have made my own classification of oxidizer in speed of reaction. Speed of reaction is not the same as strength of oxidation. E.g. persulfate is the strongest oxidizer I have, but also a rather sluggish one.

Mn2O7 (not aqueous, but very reactive)
"
and came to the conclusion that you were not talking about aqueous solutions any more so I'm not the one who moved the goalposts.

I'm not sure what colour pills I would have to be on to conclude that "aqueous solutions" and "not aqueous" were the same thing but I guess some of you could tell me.

As for "Not really, given that oxidation and reduction are two sides of the same coin." I still think its a bizare state of affairs where Li+ pops up in a list of oxidisers, technically it is one, and I dare say I could come up with circumstances where Li++ (the doubly charged ion) is an oxidant too- but it certainly wouldn't be in aqueous solution.

Make up you minds folks- if this thread is is aqueous conditions then Mn2O7 shouldn't be here and nor should Li. If it isn't about aqueous solutions then my comments about the space shuttle are legitimate.
If you are taking a rather relaxed view on the presence or absence of water, then both comments still stand.

Once you sort that out you might want to think a bit harder about who should be taking the pills.


[Edited on 29-11-2005 by unionised]
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[*] posted on 29-11-2005 at 11:46


But... but... how, if you read non-aqeuous but very reactive next to dimanganese heptoxide did you then read the rest of the list as though all the reagents were anhydrous even though some specifically say "(sluggish in aqueous media, works only in acidic media)" Perchloric acid is specifically stated at a concentration of 60% so you couldn't be thinking anhydrous and you probably know from experience the stability of the perchlorate anion in aqueous solutions. To me it was completely obvious what Woelen ment by his listing although in retrospect looking at it from your vantage point I could see where the Mn<sub>2</sub>O<sub>7</sub> could complicate things. But I believe there were enough other pointers in his post that pointed to an aqueous enviorment that would lead many to conclude they were aqueous reactions over non-aqueous reactions.

Thank's Woelen for that practical list by the way!




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[*] posted on 29-11-2005 at 13:21


I must agree with what unionised says, since redox, like many other types of reactions, are highly dependant upon conditions such as gasous or condensed phas, solvent (water or other) temp and concentrations.

my example on Li+ as oxidizer further strenthen the above mentioned. I speculate that in uncoordinative solvent the naked Li+ ion might be less stabilized and therefore more prone to reduction.

nevertheless, being right does'nt mean you can be rude. ("call NASA etc.";)
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[*] posted on 30-11-2005 at 12:50


I read the comments about non aqueous solutions and aqueous ones in the same post as evidence that people were not exclusively dealing with aqueous (or, indeed, non aqueous) reactions.
What else would I have done?
OK, the joke about 'phoning nasa might have been seen as a bit harsh, but, when it comes down to it I wasn't the first to talk about non aqueous systems and I was using it to illustrate the fact that reaction rates are so vastly dependent on circumstances that compiling a table of them is difficult, if not pointless.

On the subject of being rude I'm interested in how else (other than as being rude) one might interpret the "joke" about forgetting the pills.

The fundamental point I made - that a table of redox potentials gives both the reactant and the product, in addition to defining the conditions whereas a "speed of reactions" does none of these and will thereore be of limited value doesn't seem to have been addressed.

Of course, I look forward to a reply that addresses this tricky aspect of a ranking by rate of reaction.

[Edited on 30-11-2005 by unionised]
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[*] posted on 1-12-2005 at 12:06


I'll not waste my time any more on the issue of who's rude and what is a joke and what is a "joke". and on the subject:

"table of redox potentials gives both the reactant and the product, in addition to defining the conditions whereas a "speed of reactions" does none of these and will thereore be of limited value"

as I see it "strength" is dependent on:
1. if the reaction is termodinamically spontenious at all under the given conditions (i.e negative delta G)
2. even if we have negative delta G kinetics also play significant role. if only redox potential matter (which are only dependent on delta G since delta G = -nFE) then the wooden chair you are sitting on will start to burn since combustion of wood is termodinamically favoured... but the activation energy prevents it.
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[*] posted on 17-12-2005 at 07:59


Thanks woelen , that is a wonderful list : i have looked for such a list for a long time . It gives me several ideas as to how some natural reactions occure . With that being said I would expect that one of the titanium salts would be near the top as a very strong oxidizers ( that is just from observations of where titanium is found in natural mineralogical reactions).
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[*] posted on 9-7-2015 at 18:42


My list oxidizers:
fluorine-based
KrF+ > NiF3+ > BrF6+ > AgF2+ > PtF6 > ClF6+
oxygen-based
XeO6(4-) > O3 > S2O8(2-) > Cu2O3 > NpO3 > AmO3 > FeO4(2-) > BrO4(-) > XeO3 > Bi2O5
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[*] posted on 12-7-2015 at 04:37


Quote: Originally posted by neophyte  
Thanks woelen , that is a wonderful list : i have looked for such a list for a long time . It gives me several ideas as to how some natural reactions occure . With that being said I would expect that one of the titanium salts would be near the top as a very strong oxidizers ( that is just from observations of where titanium is found in natural mineralogical reactions).


Actually, in the presence of sunlight and H2O, I recall reading that TiO2 is a source of hydroxyl radicals!

Now, one albeit bias source (they are selling a product generating the OH radical, see http://www.hydrogenlink.com/hydroxylradicalsreactivity ) actually describes the hydroxyl radical as the ultimate oxidation tool!

So, I do agree with Neophyte comment.:)

Source: See, one of many for example: http://pubs.acs.org/doi/abs/10.1021/jp9505800

[Edited on 12-7-2015 by AJKOER]
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[*] posted on 12-7-2015 at 05:38


It seems unlikely that neophyte will see your reply nearly 10 years after they posted .
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[*] posted on 12-7-2015 at 08:45


Quote: Originally posted by unionised  
It seems unlikely that neophyte will see your reply nearly 10 years after they posted .


I guess, we will see.

The reference I cited was actual published in 1996, but internet access was probably latter, so I may have a poor excuse on being late.:P

I should also note that nano particles of TiO2 are particularly active. See, for example, "Hydroxyl radicals (.OH) are associated with titanium dioxide (TiO2) nanoparticle-induced cytotoxicity and oxidative DNA damage in fish cells", with abstract available at http://www.sciencedirect.com/science/article/pii/S0027510707... .

I did find his mineralogical viewpoint on possibly assessing chemical reactivity interesting.
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[*] posted on 12-7-2015 at 20:32


Quote: Originally posted by shadeT  
is there any safe synthesis of Mn2O7 ? where can i find information about it ?


I don't know if anyone has responded to you yet, if they have I'm sorry for repeating what they probably said.

Mn2O7 can easily be prepared as long as you have KMnO4 and 98% H2SO4.
Just add enough of the acid to cover the permanganate and you will get the oily green liquid of Mn2O7. It will be difficult to purify and will contain KMnO4 and H2SO4 impurities. D
An important note is NOT TO STROE THIS COMPOUND. Just make it when you need it. It is nasty stuff, instantly ignites room temperaure organics on fire on contact, explodes with sulfur on contact, burns plastics...nasty stuff




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[*] posted on 12-7-2015 at 20:45


Can anyone check ignite Mn2O7 pyridine, acetonitrile or not?

Sorry for my english.
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[*] posted on 12-7-2015 at 23:11


Quote: Originally posted by chemister2015  

oxygen-based
XeO6(4-) > O3 > S2O8(2-) > Cu2O3 > NpO3 > AmO3 > FeO4(2-) > BrO4(-) > XeO3 > Bi2O5


Where does one put Cu2O3, ferrates and perbromates in the topic's main table?
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[*] posted on 13-7-2015 at 06:40


Quote: Originally posted by chemister2015  
Can anyone check ignite Mn2O7 pyridine, acetonitrile or not?

Sorry for my english.


It *will* ignite dichloromethane, which is generally considered less flammable than those.




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