Is there something that can oxidise Fluorine?

[RyanJ]

Hi there everyone!

 

I've been learning about electronegativity in school and was told by my teacher that there is nothign that can oxadise Fluorine. Is this true? If not then can you tell me what the substance is that can oxadise it?

 

Cheers,

 

Ryan Jones

[BigMoosie]

I don't know of any but I wouldn't be surprised if there are, teachers often simplify things to make the syllabus easier to learn, for instance I was told none of the noble gases could form compounds but apparently there are many discovered.

Sorry not to answer your question but I just had to correct your spelling of "oxadise", it is actually "oxidise".

[RyanJ]

I don't know of any but I wouldn't be surprised if there are' date=' teachers often simplify things to make the syllabus easier to learn, for instance I was told none of the noble gases could form compounds but apparently there are many discovered.

Sorry not to answer your question but I just had to correct your spelling of "ox[b']a[/b]dise", it is actually "oxidise".

 

OOps missed that. My spelling is bad at the best of times.

 

I know teachers like to make things easier to understand but I asked this question as an extension so you'd think she would have answered correctly as far as she knew

 

I know there are things that can oxidise some of the noble gasses and thus form compounds with some of them (Fluorine!) so I guess there must be a compound that can do it to Fluorine - I'd just like to know what it is if there is one

 

Cheers,

 

Ryan Jones

[woelen]

I'm quite sure your teacher is right in this case. There indeed is no known compound, capable of oxidizing fluorine. Fluorine only exists as element in oxidation state 0, or as fluoride in oxidation state -1.

[RyanJ]

I'm quite sure your teacher is right in this case. There indeed is no known compound, capable of oxidizing fluorine. Fluorine only exists as element in oxidation state 0, or as fluoride in oxidation state -1.

 

WOW! Thats impressive!

 

I thought there would have been something that could do this but then again Fluorine has an electronegativity of 4 so unless you can get something with a electronegativit greater than that you can't.

 

If anyone else has any ideas please feel free to post but I'm guessing that woelen is probably right

 

Cheers,

 

Ryan Jones

[kortvez]

WOW! Thats impressive!

 

I thought there would have been something that could do this but then again Fluorine has an electronegativity of 4 so unless you can get something with a electronegativit greater than that you can't.

 

If anyone else has any ideas please feel free to post but I'm guessing that woelen is probably right

 

Cheers' date='

 

Ryan Jones[/quote']

 

Indeed he is right... But at the same time there is electric current which can be used to oxidize fluoride ions to elemental fluorine.

 

Zsolt

[budullewraagh]

well, define "oxidize." i consider chlorine trifluoride, for example, to be covalent, so chlorine really isn't "oxidized" per se. plus there exist compounds such as FMnO4 and FNO3, which don't exist as F+ and MnO4- or NO3-; they're more covalent, so fluorine isn't "oxidized" per se, nor is it "reduced."

[jdurg]

Oxidation doesn't require non-covalent species. If something is 'oxidized' it just means that its oxidation state has increased. So if the oxidation state of oxygen in a compound is -2, but it then it reacts and the oxidation state goes to -1, the oxygen has been oxidized even if it's not an ionic compound. You do not need an ionic species to have oxidation or reduction going on.

 

Also, your formula for FNO3 is incorrect. It should be written as NO3F, or NO2OF according to Linus Pauling. (You can see his research on the compound here.

[woelen]

well, define "oxidize." i consider chlorine trifluoride, for example, to be covalent, so chlorine really isn't "oxidized" per se. plus there exist compounds such as FMnO4 and FNO3, which don't exist as F+ and MnO4- or NO3-; they're more covalent, so fluorine isn't "oxidized" per se, nor is it "reduced."

Indeed there is a compound NO3F, but not a compound MnO4F. The highest oxidation state in which manganese exists is +7 and that is in MnO4(-), Mn2O7 and MnO3F.

 

NO3F indeed is, as Jdurg mentioned, O2NOF, with F having oxidation state -1, the middle O having oxidation state +2 and the left two O's having oxidation state -2 and the N having oxidation state -3. This is a remarkable compound with oxygen strongly oxidized by the fluorine.

[budullewraagh]

"Oxidation doesn't require non-covalent species. If something is 'oxidized' it just means that its oxidation state has increased. So if the oxidation state of oxygen in a compound is -2, but it then it reacts and the oxidation state goes to -1, the oxygen has been oxidized even if it's not an ionic compound. You do not need an ionic species to have oxidation or reduction going on."

 

see, this is a matter of personal preference. i like to say that chlorine is trivalent in chlorine trifluoride and there is a delta+ on the chlorine and delta- on each fluorine.

 

"Also, your formula for FNO3 is incorrect. It should be written as NO3F, or NO2OF according to Linus Pauling. (You can see his research on the compound here."

actually, my formula is corect. and the compound has a number of names, including: fluorine nitrate, nitryl hypofluorite, nitroxy fluoride. it's prepared by the action of fluorine on nitric acid.

 

"Indeed there is a compound NO3F, but not a compound MnO4F. The highest oxidation state in which manganese exists is +7 and that is in MnO4(-), Mn2O7 and MnO3F."

in FNO3 there's an F-O bond, last i checked. same thing in permanganyl fluoride/fluorine permanganate/permanganyl hypofluorite, which, in fact, exists

[jdurg]

"Also' date=' your formula for FNO3 is incorrect. It should be written as NO3F, or NO2OF according to Linus Pauling. (You can see his research on the compound here."

actually, my formula is corect. and the compound has a number of names, including: fluorine nitrate, nitryl hypofluorite, nitroxy fluoride. it's prepared by the action of fluorine on nitric acid.

[/quote']

 

Actually budullewraagh, you're still wrong. When naming covalent compounds, you ALWAYS name the more electropositive species first, then you name the more electronegative species. Fluorine is always the more electronegative species. Also, hypofluorite is the name of the OF- anion. If you say fluorine nitrate you're wrong because fluorine is not the more electropositive species there. Nitryl hypofluorite is wrong as well because hypofluorite is the polyatomic ion OF-. We don't call HF bifluoride because it's not an ion, so why would we all OF hypofluorite when it's not an ion? (And with NO2OF being a gas, we know that it's not ionic). Now nitroxy fluoride is a valid term because the nitroxy name would define a non-ionic NO3 group and covalent molecule nomenclature has you placing the more electronegative species last. So I will agree that nitroxy fluoride is a valid name, but fluorine nitrate is simply not correct and neither is nitryl hypofluorite. If you want to be correct about things in chemistry, then I cannot understand why you would say that the blatanly wrong naming of a compound is acceptable.

 

It's like when people say 'dihydrogen monoxide' is the chemical name for water. Well, it's not the 'correct' one (We don't go around calling H2S 'dihyrogen monosulfide', do we?). It's accepted, but the proper name if you go by the scientific rules is simply hydrogen oxide.

 

In a chemical formula, you ALWAYS put the more electropositive species first. The only exception that I know of is ammonia (NH3). So the proper name is nitroxy fluoride and the proper formula is NO3F. FNO3 would be like writing OH2 for water.

[YT2095]

I seem to remember reading somewhere that Ozone is next in line after Fluorine for Oxdation potential, can anyone confirm this at all?

[woelen]

I seem to remember reading somewhere that Ozone is next in line after Fluorine for Oxdation potential, can anyone confirm this at all?

No, there are other stronger oxidizing agents, such as PtF6, which is capable of oxydizing oxygen. Another really strong oxidizer is sodium perxenate, Na4XeO6.8H2O. But indeed, ozone is a very strong oxidizer, similar in strength as peroxodisulfate (redox potential just over 2 volts).

 

@budullewraagh: Permanganyl fluoride is not MnO4F, but MnO3F, with a direct Mn-F bond and Mn in its +7 oxidation state. I recently discovered that I actually made some of this stuff in one of my home experiments, it is a green volatile compound. Just google on permanganyl fluoride and you'll see a few links, where the formula MnO3F is given to this compound. I think that MnO4F does not exist.

 

You can see this green stuff in the following link:

http://woelen.scheikunde.net/science/chem/exps/KMnO4+NaF+H2SO4/index.html

 

When I did the experiment I did not know it, but some literature study has shown me that the green compound is MnO3F.

[RyanJ]

I seem to remember reading somewhere that Ozone is next in line after Fluorine for Oxdation potential, can anyone confirm this at all?

 

Isn't Ozones oxidation potential 2.07V? If it is then its lower than that of oxygen which has an oxidation potential of 2.5V.

 

Cheers,

 

Ryan Jones

[woelen]

Isn't Ozones oxidation potential 2.07V? If it is then its lower than that of oxygen which has an oxidation potential of 2.5V.

 

Cheers' date='

 

Ryan Jones[/quote']

 

Are you sure about the oxygen potential of 2.5 V? Even 2 volts already is very strongly oxidizing and the only readily available chemicals which are just over this are salts of peroxodisulfate and ozone.

 

Have a look at this table:

 

http://members.aol.com/logan20/red_tabl.html

 

Here oxygen, acting as oxidizer in acidic environment has a redox potential of 1.23 volts.

[RyanJ]

Are you sure about the oxygen potential of 2.5 V? Even 2 volts already is very strongly oxidizing and the only readily available chemicals which are just over this are salts of peroxodisulfate and ozone.

 

Have a look at this table:

 

http://members.aol.com/logan20/red_tabl.html

 

Here oxygen' date=' acting as oxidizer in acidic environment has a redox potential of 1.23 volts.[/quote']

 

Just goes to proove that the internet is not always the best source of information

 

I'll get the CRC book out and see if it has it in there - give me a few minutes:

 

Here we go:

 

Oxygen: [ce]O2 + 4H^{+} + 4e^{-} <=> 2H2O[/ce] 1.229

Ozone: [ce]O3 + 2H^{+} + 2e^{-} <=> H2O + 02[/ce] 2.076

 

Cheers,

 

Ryan Jones

[YT2095]

I was thinking/speaking in terms of a stand alone element rather than a compound in this instance.

 

and it`s much Higher than just plain O2, I`m pretty sure of that

[verode]

YES you may get Fluoriine

SbF5 is a great Lewis Acid

KMnF5 es stable but the mix

get SbF6(-1) and MnF4 instable

then MnF4=MnF3+ fluorine

 

NO3F is a trong oxidant an you need fluorine

Ther are very interisting reactions

like AgF with bromine and you get BrF3

or iodine you get IF5 very strongs oxiders

[insane_alien]

YES you may get Fluoriine

SbF5 is a great Lewis Acid

KMnF5 es stable but the mix

get SbF6(-1) and MnF4 instable

then MnF4=MnF3+ fluorine

 

NO3F is a trong oxidant an you need fluorine

Ther are very interisting reactions

like AgF with bromine and you get BrF3

or iodine you get IF5 very strongs oxiders

 

and in not a single one of those is fluorine oxidised.

[verode]

the more powerful oxidant is the ion KrF+ or some high oxidations compound of nikel

you get thanks Lewis acid's like SbF5+KrF2

[DrP]

I know there are things that can oxidise some of the noble gasses and thus form compounds with some of them (Fluorine!) so I guess there must be a compound that can do it to Fluorine...

 

The compounds which oxidise the noble gasses mainly contain Fluorine iirc. XenonFluride (XeF6)is a good example of a possible xenon compound.......and guess what?? It's the fluorine doing the oxidising!

 

Woelen posted about Na4XeO6.8H2O being a really strong oxidizer - this is probably due to the fact that the Xe is so weak at keeping it's oxygen that it readily gives it up (it's stable enough without it anyway).

[John Cuthber]

And fluorine still remains unoxidised.

Just a thought, if you are going to raise a thread like this from the dead, you might want to check that what you are saying is actually worthwhile.

[nitric]

fluorine is set up to be unoxidizable by existing in the 1+ and the 0 ox. states

[John Cuthber]

Is there a term for drive-by necromancy?

Anyway I can oxidise Fluorine perfectly well in an electric arc. I just can't get the F++ ions to form a compound with anything

[nitric]

no it oxidixzes everything in an electric arc, like when they make [ce]F_{2}O[/ce] in a lab(though its pretty hard)or thats what i belive, since i learn that theres only two ox. states to fluorine(1+ and 0)