In oxidation-reduction or redox reactions, it is important to be able to identify which atoms are being oxidized and which atoms are being reduced. To identify if an atom is either oxidized or reduced, you only have to follow the electrons in the reaction.
Identify the atoms that were oxidized and which atoms were reduced in the following reaction:
+ 2 Al → Al2
+ 2 Fe
The first step is to assign oxidation numbers to each atom in the reaction. The oxidation number of an atom is the number of unpaired electrons available for reactions.
Review: Rules for Assigning Oxidation Numbers
The oxidation number of an oxygen atom is -2. 3 oxygen atoms has a total charge of -6. To balance this, the total charge of the iron atoms must be +6. Since there are two iron atoms, each iron must be in the +3 oxidation state. To summarize: -2 electrons per oxygen atom, +3 electrons for each iron atom.
The oxidation number of a free element is always zero.
Using the same rules for Fe2
, we can see there are -2 electrons for each oxygen atom and +3 electrons for each aluminum atom.
Again, the oxidation number of a free element is always zero.
Put all this together in the reaction, and we can see where the electrons went:
Iron went from Fe3+
on the left side of the reaction to Fe0
on the right. Each iron atom gained 3 electrons in the reaction.
Aluminum went from Al0
on the left to Al3+
on the right. Each aluminum atom lost three electrons.
Oxygen stayed the same on both sides.
With this information, we can tell which atom was oxidized and which atom was reduced. There are two mnemonics to remember which reaction is oxidation and which reaction is reductions. The first one is OIL RIG
oss of electrons
ain of electrons.
The second is "LEO the lion says GER".
lectrons in O
lectrons in R
Back to our case: Iron gained electrons so iron was oxidized. Aluminum lost electrons so aluminum was reduced.