Definition 4

So far, we have gone through 3 definitions of oxidation. But this last definition, is actually the most useful and popular.

The oxidation number/state of an atom in a compound is equal to its theoretical charge if all the atoms were joined together by electrovalent bonds.

Thus, an increase in oxidation number means the substance has been oxidised.

Before we can apply this definition, there are 4 basic rules to follow:

Rule	Example	Oxidation State
The oxidation state of a free element is zero	Cu S Cl2	0 0 0
The oxidation state of a simple ion is the same as the charge on the ion	K+ Al3+ S2-	+1 +3 -2
The oxidation states of the atoms present in the formula of a compound add up to zero.	CaCO3 [Ca2+; CO32-]	+2 – 2 = 0
The total of the oxidation states of the atoms in a polyatomic ion is equal to the charge on the ion.	SO42-	+6 + 4(-2) = -2

Remember, when writing the oxidation state, the charge comes first than the charge value.
for example, the charge of an oxide ion is 2-, but its oxidation state is -2.

You can read more about the rules of oxidation states. click on this link.
http://www.chemtutor.com/redox.htm#state



After you have viewed the video, let's apply the concept of oxidation numbers to chemical reactions.

Example 1
Fe2+ + H2O2 → Fe3+ + H2O
In this reaction, has Fe2+ been oxidised or reduced?
the oxidation number of Fe2+ has been increased from +2 in Fe2+ to +3 in Fe3+.
Thus, Fe2+ has been oxidised. The oxidising agent is H2O2.

Example 2
3SO2 + Cr2O72- +2H+® 3SO42- + 2Cr3+ + H2O
In this reaction, has dichromate(VI) been oxidised or reduced?
the oxidation number of Cr in Cr2O72- has been decreased from +6 in Cr2O72- to +3 in Cr3+.
Thus, Cr2O72- has been reduced. The reducing agent is SO2.



After understand all the four definitions of oxidation, we will move on to learn more about the oxidising and reducing agents.