Ah. Oxidation numbers. How delightful.
I think of it this way. The oxidation number of an atom in a compound is the average charge the atom would have if all bonds were considered to be fully ionic. In covalent bonds, the electrons are treated to belong to the more electronegative atom.
In a truly equal bond, like the nitrogen triple bond with itself, the electrons are really shared and the oxidation number is, well, zero.
In something like water, as hydrogen is less electronegative, both hydrogens are treated as not owning the electrons in the covalent bonds, so the oxidation number of hydrogen there is +1 (each hydrogen is seen as losing one electron each for a total charge of +2 for two hydrogens; average is +1). Similarly, the oxygen gains an electron from each hydrogen by this model => oxidation number of -2. Indeed, as oxygen is so electronegative, we normally assign its oxidation number automatically as -2 and fluorine's as -1 (true for ALL fluroine compounds except F2). We assign hydrogen as +1 automatically because its electronegativity is quite low and it tends to lo se its electron.
What about compounds that have fractional oxidation numbers? Let's consider something like propane. Using conventional methods, we find that propane has an oxidation number on the carbon as -8/3. This means that the AVERAGE of the individual oxidation numbers if -8/3. Let's see why.
Propane can be written as CH3-CH2-CH3. The first and third carbons have three hydrogens attached each, so they gain three electrons each (remembering that carbon is more electronegative than hydrogen but only marginally). They are both bonded to the second carbon as well but C-C bonds are symmetrical => no contribution to O.N here.
Now, the second carbon only has two hydrogens attached, so its oxidation number is -2. We have two -3s and a -2 for an average of -8/3. This can be extended to more complex molecules, but it's not really needed for the VCE course. Indeed, you can get by in VCE without knowing what oxidation numbers really are, but I think it's helpful if you had an idea of where they came from.
Reduction = oxidation number decreases. Reductant causes the other reactant to be reduced (it itself is oxidised). Work out the others yourself from this.
O has an oxidation number of -2 (in most cases). Therefore, Iron on the left has an oxidation number of -3 and an oxidation number of 0 on the right. Carbon on the left has an oxidation number of -2, and on the right an oxidation number of -4.
You can do the rest now 
Typos; iron oxidation number is +3
also, the carbon ONs are both positive as they're both bonded to oxygens.
Metals almost always have positive ONs and their ONs are equal to their charges if they exist as ions.
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