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Strenght of oxidizers

vulture - 6-9-2002 at 07:02

As we frequently use several oxidizers, react them with eachother and synthesize them, I thought it would be useful to compile a table of strength of oxidizers.

First of all, there are a few classification problems. We could just take the redox potential and sort according to that, but that doesn't always correctly predict behaviour and also is rather confusing to work with (mmkay, this get's oxidized then that is reduced and..err. oh..or is this the oxidizer?).

Secondly, the reaction situation has to be determined. I was thinking of the strength in solution, simply by testing which oxidizer would be capable of oxidizing the other.

These are the compounds I'd like to see classified (I've classified them according to what I think, starting with the strongest):

Mn2O7
HClO4
S2O8 2-
Cr2O7 2-
MnO4 2-
ClO2 -
ClO -
ClO3 -
H2O2
ClO4 -

Remember that this is all in solution regarding cold chemical oxidation.

Waiting for your corrections/comments...

vulture - 6-9-2002 at 07:03

The H2O2 should be listed above ClO3-

Polverone - 6-9-2002 at 14:29

You're right that different oxidizers have different behavior depending on conditions. H2O2 can serve as a reducing agent or an oxidizer, for example. We're often interested in oxidizers in high-temperature anhydrous conditions (combustion) as well as aqueous solution. Ammonium perchlorate is a much better high-temperature oxidizer than sodium hypochlorite, but sodium hypochlorite will react with many more chemicals in solution.

I think that I would add HNO3 to your list somewhere in the upper half.

I'd also add a few mild oxidizers: Cu2+, Pb4+, and MnO2. They would probably fall near the bottom of the list.

As far as your ordering goes, I think that persulfate salts are considerably less aggressive than permanganates and dichromates in aqueous solution. Also I think that chlorates are far less active than H2O2.

There are also perchromates and manganates, which are powerful oxidizers (the perchromate especially).

Somewhere on that list should go the halogens, although none of us is very likely to be using fluorine or astatine in the lab.

vulture - 8-9-2002 at 01:35

persulfates are very strong oxidizers, when moist, they will spontaneously decompose into O2 with a high O3 percentage. Also, it is known that persulfate will oxidize chlorate to perchlorate. My attempts with MnO4 - to do this have failed so far.

shadeT - 13-10-2003 at 09:42

is there any safe synthesis of Mn2O7 ? where can i find information about it ?

BromicAcid - 13-10-2003 at 13:09

Ive heard that BiO3- is a plenty powerful oxidizing agent. My inorganic text book has this to say

"One of the few oxidizing agents even more powerful then permanganate is the bismuthate ion [BiO3]- A test for manganese (II) ion is the addition of sodium bismuthate to a sample under cold, acidic conditions. The purple permanganate ion is produced, thereby indicating the presence of manganese."

Descriptive Inorganic Chemistry [Third Edition] Geoff Rayner-Canham Tina Overton

chemoleo - 13-10-2003 at 15:45

thats interesting (the BiO3- bit). Although it's off topic, how would I go about making it? Assuming my starting point is Bi metal?

JustMe - 14-10-2003 at 17:42

According to my old (1961), but very useful Textbook of Inorganic Chemistry (Partington)... On fusing Bismuth Trioxide with Potassium Hydroxide in air, a brown mass of Potassium Bismuthate, KBiO3, is formed. It is hydrolyzed by water, but with cold solutions of Manganous salts in dilute nitric acid it gives a purple solution of Permanganic Acid. So there you go.

Bismuth Trioxide, Bi2O3 is formed by heating the metal, hydroxide, basic carbonate or nitrate to redness in air. Hmm, just noticed... it also says that Bismuth Trioxide is used to produce an iridescent white glaze on porcelain. Perhaps it is available through a pottery supply company???

Not mentioned in your list, but leave us not forget the supreme oxidizing power of Xenon Difluoride!

bismuth

chloric1 - 4-1-2004 at 13:40

You know with an ebay and all Bismuth is readily available and it is not a hazardous material. And yes it is tricky on figuring outoxidizer strength. ClO3 is stronger than H2O2 in the respect that ClO2 oxidizes H2O2 to clorous acid and oxygen gas; but peroxysulfates can oxidize chlorates to perchlorates. I always regarded peroxysulfates as hydrogen peroxide adducts. Confusing indeed.

BromicAcid - 4-1-2004 at 14:35

Quote:

Not mentioned in your list, but leave us not forget the supreme oxidizing power of Xenon Difluoride!

I always thought that the oxidizing power of Xenon Difluoride came from it releasing free fluorine in the reaction mixture. For the list, at the very top, I would just put F2 fluorine has been described as the "Tyrannosaurus Rex" of the periodic table. I mean, how else am I going to make perbromic acid?

unionised - 4-1-2004 at 15:35

IIRC CoF3 decomposes to F2 and CoF2 on heating. This means that Co (III) is a stronger oxidant than F2
(Yes, I know that's blasphemous).
OTOH Co(OH)2 is oxidised by air to give Co(III) so O2 is a better oxidant than Co(III)
This means that O2 must be a stronger oxidant than F2. The trouble is that we know it isn't.
This serves to ilustrate the fact that what is a good oxidant depends on circumstances to such a great extent that putting them in order is going to be a bit of a non starter unless you use some sort of standard conditions. If you do that, you might as well use the electode potentials.
Add to this the problem of kinetic versus thermodynamic stabillity (Fe(II) perchlorate shouldn't exist) and the list won't really help very much.
BTW if you react peroxide with acetic anhydride you get peracetic acid. This will oxidise Mn(II) to permanganate, even in slightly acid conditions.
(Peracetic acid is rather unstable and I wouldn't use it if I were you)

blip - 7-1-2004 at 22:25

Quote:

Add the H2O2 to the beaker(or drop reactant into H2O2) and test for gas bubbles.
If there are any trap in a test tube and test for any explosive peroxides using a lighted splint.

Are you saying that the organic peroxides will certainly be in the gas phase with any oxygen produced? I don't think so... oxygen is the gas produced from H2O2 decomposition (assuming that's all that's occurring) and it lights a smoldering splint. I don't think that would be a good test at all considering the common production of oxygen would cause a false positive. Did you mean to say something completely different and typed out the wrong thought?

BromicAcid - 27-11-2005 at 13:05

I was looking around the internet for a table listing the electrochemical series and found one that has a number of enteries in that that would be considered extrenuous on a normal table, this one has 536 enteries on it and the top ones are whoppers such as:

XeF + e- = Xe + F- 3.4V
Cf4+ + e- = Cf3+ 3.3V

And many more interesting enteries (ferrates, fluorine/oxygen compounds...), see what I mean at:

http://www.efunda.com/materials/corrosion/electrochem_list.c...

[Edited on 11/27/2005 by BromicAcid]

The_Davster - 27-11-2005 at 13:31

WOW Bromic

. I have been looking for such a giant table for a long time. Thank you.

Pr4+, Tb4+ !! I gotta get me some more lanthanides

EDIT: AArg, I got through perhaps the first 200, and it wants me to subscribe :mad:

[Edited on 27-11-2005 by rogue chemist]

Phel - 27-11-2005 at 14:14

Quote:

EDIT: AArg, I got through perhaps the first 200, and it wants me to subscribe :mad:

Yes, I was looking forward to see some exotic reducing agents when I recieved the same thing. I will add the electrochemical series listed in the CRC to Madhatters FTP when I find it. Same content, but as a PDF.

Edit: (Madhatters FTP) Upload -> CRC_SRPotentials.pdf

[Edited on 27-11-2005 by Phel]

Pommie - 27-11-2005 at 15:35

Quote:

AArg, I got through perhaps the first 200, and it wants me to subscribe :mad:

When you get this problem, go to BugMeNot and enter the site name. For this site (efunda) it comes up with:-
Username mrx
Password xyz

Mike.

[Edited on 27-11-2005 by Pommie]

The_Davster - 27-11-2005 at 16:32

I had tried that already, aparently mrx dident pay for this years subscription. Oh well, you have access again in around an hour and by modifing the url you can get to the later pages. I have got to around 300 saved now.

nitroglycol - 27-11-2005 at 16:34

Hmm. Interesting that that site lists oxygen difluoride below fluorine; I'd expect it to be higher.

Dr. Beaker - 28-11-2005 at 09:17

I think a powerfull, possibly the strongest oxidizer might be KrF2 since the Kr-F bonds are even weaker then in elemental F2 (and are among the weakest known covalent bonds to exist).
weaker, less "exotic" (and more safe) oxidizer are perxenates: XeO4-2 (XeO3 and XeO4 are explossivly unstable). the big advantage of using them is that the byproduct is Xe, hence no need for seperation and purification of products (like H2O2 but much stronger)

Dr. Beaker - 28-11-2005 at 09:30

and some more on the subject...
some time ago I heard a seminar in our department and what catched my attention (between snors...) was a certain lithium complex of an organic molecule that undergo disproportionation to...
Lithium metal

and a some other specie
the lithium precipitate as a shining miror on the flask's wall...
i.e - this molecule is a stronger reducer then metallic Li.

unionised - 28-11-2005 at 11:51

" this molecule is a stronger reducer then metallic Li."
So, it's about as far off-topic as you can get.

woelen - 28-11-2005 at 12:39

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)
MnO4(-)
ClO(-)
Cl2
HNO3 concentrated
H2O2
Cr2O7(2-)
Br2
H2SO4 concentrated
S2O8(2-) (sluggish but potent in aqueous media)
BrO3(-) (sluggish in aqueous media, works only in acidic media)
I2
ClO3(-) (sluggish in aqueous media, works only in acidic media)

HClO4 (hardly works as oxidizer, not at 60%, just a nice clean strong acid)
ClO4(-) (not oxidizing at all, totally inert)

So, my list is not an electrochemical series, but more like a series of observed reaction speeds. Mn2O7 is very fast acting and HClO4 (up to 60%) and ClO4(-) are so slow, that they cannot be considered oxidizers anymore.

I took some 60% HClO4 and added solid KI. This does not give any reaction (besides dissolving some KI). Even when heated close to boiling, the liquid only becomes a little yellow/brown, that's all.

With a mix of HClO4 and NaCl, only when heated to boiling, a very faint smell of chlorine can be observed. From these observations I conclude that HClO4 (at least not at 60 .. 70%) is not useful as oxidizer at all, despite all the strong stories on the Internet about the dangers of HClO4.

unionised - 28-11-2005 at 12:47

Phone NASA and tell them to cancel the next shuttle- perchlorates aren't fast oxidisers any more.

There is probably less hope of getting a valid list of oxidisers (fast to slow) than there was of getting one (strong to weak).
At best you would need the pre- exponential and the activation energy terms for an Ahrenius type equation. I think that means you would need a veactor for each oxidant and I'm not sure how you could put those in order.

Dr. Beaker - 28-11-2005 at 13:08

unionised,
what is your problem?
have you mixed your medications today?
have'nt the nurse told you to take the blue pills?

nitroglycol - 28-11-2005 at 16:11

Quote:

Originally posted by unionised
" this molecule is a stronger reducer then metallic Li."
So, it's about as far off-topic as you can get.

Not really, given that oxidation and reduction are two sides of the same coin.

woelen - 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.

unionised - 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]

BromicAcid - 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 Mn2O7 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!

Dr. Beaker - 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."

unionised - 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]

Dr. Beaker - 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.

neophyte - 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).

chemister2015 - 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

AJKOER - 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]

unionised - 12-7-2015 at 05:38

It seems unlikely that neophyte will see your reply nearly 10 years after they posted .

AJKOER - 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.

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.

The Chemistry Shack - 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

chemister2015 - 12-7-2015 at 20:45

Can anyone check ignite Mn2O7 pyridine, acetonitrile or not?

Sorry for my english.

ave369 - 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?

DraconicAcid - 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.