[WhiteRose]

Hey guys, I was wondering if anyone can help me.
Which one of the following are oxidizing agents?
KI, K2S2O8,  Na2S2O3
Thanks

[cuongt]

well i know that KI is one but dont know the others

[hmx9123]

I would not consider KI an oxidizing agent at all.  It is a salt.  What you need to look for in an oxidizing agent is an element at an unusually high oxidation number--i.e., it wants electrons, and thus can 'oxidize' other materials by taking their electrons.  For example, the I^- in KI is -1, which is normal for iodine.  If it were KIO₄, then you have I in the +7 oxidation state, thus wanting electrons desperately and being a powerful oxidizing agent.

Think about the oxidation states of sulfur in K₂S₂O₈ and Na₂S₂O₃.  Compare them to one another and to sulfur's normal oxidation state.  What is sulfur's normal oxidation state? (you may have to look at the periodic table to figure this out).  Which ever one has a higher oxidation number will be your more oxidizing compound.  If you have trouble calculating oxidation numbers, post your attempts at it, and we will help you work through it.

[Borek]

What you need to look for in an oxidizing agent is an element at an unusually high oxidation number--i.e., it wants electrons, and thus can 'oxidize' other materials by taking their electrons.

Do you consider chlorine to be oxidizing agent, or not? Id doesn't fit the idea of 'unusually high oxidation number'.

Think about the oxidation states of sulfur in K₂S₂O₈ and Na₂S₂O₃

Very misleading approach. K₂S₂O₈ properties are no sulfur dependent (both S being +6, just like in SO₄^2- -- which is hardly powerfull oxidizer). Na₂S₂O₃ has two completely different sulfurs and its properties are determined mostly by the sulfur with -2 oxidation state.

I think you have moved a little bit too far trying to simplify

[Mitch]

Borek:
I thought H₂SO₄ was a strong oxidizer?  

Shouldn't Sulfur's oxidation number be +7 not +6 for K₂S₂O₈

and +2 not -2 for Na₂S₂O₃

[hmx9123]

Consider:

Cl₂ is an oxidizer.  Cl has an oxidation number of 0.  This is unusually high for chlorine which is usually -1.  Maybe not a great order of magnitude, but it is not the usual oxidation state of chlorine.  Therefore it is unusual.  Unusual doesn't necessairly imply a radical magnitude difference, although I could see where one could imply that from what I wrote.

Secondly:

As far as I'm aware, sulfuric acid isn't a powerful oxidizing agent; it is a powerful dehydrating agent, though.  You may be thinking of nitric acid, Mitch, which is a significant oxidizer.

Perhaps this is a bit oversimplified, but for most high school classes, these are the rules used to determine the oxidation states of the hetero atoms.

1. All elements are in the 0 oxidation state in their elemental form
2. Alkalis are always +1
3. Oxygens are usually -2, although can be -1 in peroxides and fractional in superoxides.
4. Hydrogen is usually +1 although can be -1
5. Most other elements are the charge you would expect from the periodic table, however you must achieve charge balance, and therefore these elements can be of differing oxidation states.

Now, that being said (and that was off the top of my head, so it may not be complete), there are some more complicated situations that arise when several atoms of the same type have different oxidation numbers, such as Fe₃O₄.  (Fe(II) and Fe(III) mix).

Going on what most HS students are taught:

Sulfur in sulfuric acid is in the +6 state.  
Sulfur in K₂S₂O₈ is +7 as Mitch said, and is +2 in Na₂S₂O₃ as he also stated.  

Thus, potassium peroxydisulfate is a more powerful oxidizing agent than sodium thiosulfate.

Although the situation is more complicated than that if you look at the structures of these compounds, I would not expect a HS student to understand any more than that.  If someone wants to post a structure, we can go into more detail, but the underlying idea that one of these materials is more oxidizing based on average oxidation number still works.

[AWK]

Potassium persulfate (OxN +7)  (K2S2O8) is a strong oxidizer, sodium thiosulfate has reducing properties, it can be even oxidised by elemental iodine.

Mitch - H2SO4 (OxN +6) sometimes is also an oxidizer. Hot concentrated H2SO4 dissolves copper or silver
eg:
Cu+2H2SO4=CuSO4+SO2+2H2O

[hmx9123]

AWK--that's pretty cool about the hot concentrated sulfuric acid dissolving copper.  I actually didn't know that reaction occured.

[Borek]

Oxidation numbers are dangerous and I am trying to avoid them where possible. S₂O₈^2- properties are not due to +7 of sulfur, but due to peroxy bridge (hmx: note one of your rules states that oxygen can -1!). Sulfur in thiosulfate is +2 on average, but in fact there are two different sulphurs - +6 and -2.

I have stated this several times - we teach chemistry using simplifications, but we should be very carefull to not oversimplify and to avoid using examples that are not correct. If we state "metal oxides react with water to produce bases" we can use example of Na₂O or CaO, but not Al₂O₃. If we state (as hmx did) "oxidizers are substances with atoms at unusually high oxidation numbers" we can use permanganate, chlorate or iodate as examples, but not peroxydisulfate, as its properties have nothing to do with sulfur oxidation number.

In other words - simplifications - yes, things that have to be 'unteached' later - no.

[hmx9123]

Borek--
  I do agree that we should not oversimplify things; that generally makes a mess of things.  Perhaps it is best to go through the complicated (for HS students) structural drawings and explanations.  Most likely the book itself asked poor questions of the students and in certain ways forced them to oversimplify.

I am aware that the peroxy bond in the peroxydisulfate is the source of the oxidation--and that in turn leads to a sulfur oxidation number of +6, not +7.  It also means that my original statement about substances with unusually high oxidation numbers being oxidizers still stands--the high oxidiation number being oxygen at -1.  I also stated that you want something that can take electrons away from another compound in order to oxidize.  That is a much more general rule.

Aside from this, I agree, we definitely need more examples where we can use the simple rules without teaching incorrect information.  We should probably teach more exceptions, although they are generally glossed over either due to time or in an attempt to keep students from being confused.  There is a fine line to walk, though, as teaching nothing but exceptions almost teaches the opposite trend of what needs to be taught.

To answer the question asked, if we need a basis for comparison, how is a HS student supposed to come up with the oxidizer answer without knowing this information about the structures and the exceptions specifically?  A high school student would be lucky to draw a correct lewis structure of the peroxydisulfate anion, seeing as many of them haven't seen peroxide bonds.  And, if you have sulfurs of differing oxidation states, as in the thiosulfate, how would a student be expected to compare this to the oxidizing ability of the oxygens in the peroxydisulfate anion?  The oversimplification to look at the oxidation numbers of the sulfur may be incorrect, but it is the answer that the HS texts allude to.  That is kind of a shame, really.  I, too, wish that books would use examples such as the perchlorates and chlorates rather than those which are exceptions to the general rules.

[hmx9123]

Actually, our discussion poses a good topic of chemical education discussion, and I have started another thread in that forum.  I'd like for anyone interested in the teaching of oxidation numbers to check it out.

[Borek]

That is kind of a shame, really.  I, too, wish that books would use examples such as the perchlorates and chlorates rather than those which are exceptions to the general rules.

If so - we agree I just can't force myself to give answers that are wrong - I am not going to dilute my reputation (if I have any ) supporting other's mistakes.

[hmx9123]

Well, check out the new thread and let's figure out some new ways of teaching.