oxidation states

[akcapr]

Oxidation sattes, they confuse me. What exactly are they? i was lookin at this part of woelens site V and the manganese and iron oxidation states.. -

 

(http://81.207.88.128/science/chem/exps/bleach/index.html)

 

confused me. could someone explain the oxidation states and how it haopens, and wat the products are. thx

[budullewraagh]

"oxidation is the loss of electrons. reduction is the gain of electrons." <--what you'll hear in high school chem.

 

that is true, although it's not exclusive for electrostatic ionic bonds. it works for polar covalent bonds as well.

 

consider LiF; the Li is at +1 and the F is at -1. The lithium cation has an electron configuration equivalent to that of helium and fluoride anion has an electron configuration equivalent to that of neon.

 

now consider ClO4-. the O all needs to "reduce" to "-2" (even though the Cl-O bond is polar covalent and far from "ionic.") 2*4=8. ClO4- is -7, so Cl must be at +7.

 

 

 

in other news, i knew NaOCl was a powerful oxidizer and that the anhydrous form is very explosive, but i never knew it could oxidize iron to form ferrates. i've heard of the fused KOH+KNO3+Fe and BaO2+Fe2O3, and the electrolytic methods, but recently i've been impressed by the oxidizing abilities of NaOCl. interesting, that

[latentheat]

Oxidation states indicate the number of electrons lost or gained (or the apparent loss or gain) by an atom in a compound. If something is in the oxidation state of +2 for example, like Mg in MgO, it has lost 2 electrons.

 

Also, from another prospective, if you have CuSO4 solution, the Cu ions in solution will be in the +2 oxidation state, and to precipitate copper you have to reduce it by giving it 2 electrons. This can be done by electrolysis if the metal to replace it is less reactive than copper, or spontaneously if the metal to replace it is more reactive than copper.

 

Cr in the +6 oxidation state looks orange or reddish. Cr in the +3 oxidation state is greenish. That's why changing the oxidation states can cause nice color changes.

[akcapr]

ah that helps some. but im confused about this quote from woelens site:

 

On heating, the hypochlorite is capable of oxidizing the ferric hydroxide and hydrous manganese (IV) oxide further:

 

2Fe(OH)3 + 3ClO– + 4OH– → 2FeO42- + 3Cl– + 5H2O

 

2MnO2 + 3ClO– + 2OH– → MnO4– + 3Cl– + H2O

 

can someone explain why heating causes "further oxidation" (???) and what exactly is happenning there. thx for putting up with my dumbnesseness.

[latentheat]

The iron is going from +3 to +6 in the first reaction. In the second reaction the manganese is going from +4 to +7. That's what he means by further oxidation - they are losing more electrons.

 

Heat is the activation energy for the reaction.

[budullewraagh]

woelen, on your site the reaction:

2Fe(OH)3 + 3ClO– + 4OH– → 2FeO42- + 3Cl– + 5H2O

is a bit unclear to me. surely the hydroxide must be fused, not aqueous, right?

[akcapr]

isnt it part of the iron hydroxide?

[budullewraagh]

the iron is oxidized to +3 and is the hydroxide. ferric hydroxide is completely insoluble in water.

[woelen]

woelen' date=' on your site the reaction:

2Fe(OH)3 + 3ClO– + 4OH– → 2FeO42- + 3Cl– + 5H2O

is a bit unclear to me. surely the hydroxide must be fused, not aqueous, right?[/quote']

Just plain aqueous. No fusing at all...

[woelen]

Oxidation state is just a tool for bookkeeping of electrons. Bookkeeping is done, such that it gives a good impression of how far an element is oxidized (or reduced).

 

Some elements gain electrons easily (they act as oxidizers and they are reduced). The most common is oxygen. The oxo-group (or oxide ion) is said to have oxidation state -2. Hydrogen usually has oxidation state +1 in its compounds (except in metal hydrides, where it has oxidation state -1). Using these simple rules, one can easily compute the oxidation state of other elements.

 

Example: MnO4(-). Total charge: -1. Charge per oxo-group: -2. So, you need +7 for manganese in order to get a total charge, equal to -1.

Keep in mind though, that in MnO4(-), the manganese does not really have a charge equal to +7. In many real life compounds, the charge is not distributed as extremely as oxidation numbers would suggest. So, in reality, the charge on the manganese will be just a little over 1 and the charge on each of the oxo-groups is a little below 0, such that the total charge still remains -1.

The usefulness of oxidation numbers is in that they easily tell you whether an element is strongly oxidized or not. In my experiment, you have Fe(OH)3, with Fe in the +3 oxidation state. In the ferrate ion, the iron has formal oxidation state +6. So, from this playing with numbers, you immediately see, that the ferrate is an iron-species, which is oxidized further than ferric hydroxide.

 

Sometimes the concept of oxidation number is flawed. Consider the deep blue compound CrO5. This does not contain chromium in the +10 oxidation state, but it contains chromium in the +6 oxidation state, with an oxo-group attached to it and two peroxo-groups attached to it. For one oxygen atom we have oxidation state -2 and for the other four oxygen atoms we have oxidation state -1. A better formula is CrO(O2)2. So, with oxidation numbers you have to be careful and you have to know your compounds.

 

Just another nice one: try to determine the oxidation state of S, C and N in the thiocyanate ion SCN(-) and in the thiosulfate ion S2O3(2-). Here you'll also see that the concept of oxidation number sometimes is not that easy.

[budullewraagh]

"Just plain aqueous. No fusing at all..."

 

thats very, very surprising. Fe(OH)3 is almost completely insoluble in water and i've NEVER seen KOH used as an oxidizing agent in any state except fused

[woelen]

"Just plain aqueous. No fusing at all..."

 

thats very' date=' very surprising. Fe(OH)3 is almost completely insoluble in water and i've NEVER seen KOH used as an oxidizing agent in any state except fused[/quote']

Well, I must say, I also never have seen KOH or NaOH as oxidizing agent . In fact, these are not oxidizing agents. What you are referring to is that KOH is used as supplement in order to facilitate a certain redox reaction (e.g. oxygen from the air, or KNO3 is the oxidizer and the NaOH or KOH supports the reaction).

 

Now coming back to your original question.

1) Fe(OH)3 is almost insoluble in water: This is true, but it does not prevent its oxidation. The metal zinc also is completely insoluble in water, yet it can be oxidized easily to Zn(2+) with just plain acid.

 

2) KOH (or NaOH) used as supplement to a redox reaction: This in fact is nothing special. At high schools, people are used to redox reactions, with H(+) as supporting reagent (e.g. Cr2O7(2-) or MnO4(-) as oxidizer in acidic environments), but the only reason for this is that the oxidizers, mentioned at high schools, usually contain a lot of oxo-groups, which must be neutralized with H(+). These common redox reactions use up H(+), and hence these reactions run best in acidic environments.

In the reactions, which I show in the experiment, it is the other way around. Look at the reaction equations for formation of ferrate or permanganate, and you'll see that these reactions use up OH(-), so these reactions run best in alkaline environments.

 

Some other examples of redox reactions, which require the ion OH(-) for their completion:

 

Oxidation of chromium (III), e.g. chromium hydroxide or a green solution of sodium chromite, with H2O2. In strongly alkaline aqueous solutions, you can easily oxidize chromium (III) to the yellow chromate with a little heating.

 

Oxidation of manganese (II) to manganese (IV) by oxygen from the air. This only happens in alkaline environments, because this reaction uses up OH(-).