REDOX REACTIONS
In oxidation reduction or redox reactions, as reactants form products, one or
more atoms are reduced while one or more other atoms are oxidized. If reduction
occurs in a reaction, then oxidation must also take place. If one species gains
electrons, another must have lost them. If we want to represent just the reduction
or just the oxidation in a redox reaction, we write something called a half-
reaction.
Look at this reaction.
Fe + 2HCl → FeCl 2 + H 2On the left side of the equation, iron’s oxidation state is 0 and hydrogen’s is +1.
On the right side of the equation, iron’s oxidation state is +2 and hydrogen’s is 0.
So iron has been oxidized (each atom has gone from an oxidation state of 0 to
+2, so each has lost two electrons), and hydrogen has been reduced (each atom
has gone from an oxidation state of +1 to 0, so each has gained one electron). We
can write two half-reactions, one showing the oxidation of iron and the other
showing the reduction of hydrogen. Here they are.
Oxidation: Fe → Fe+2 + 2e−
1 iron atom loses 2 electrons and takes on an oxidation state of
+2.
Reduction: 2H+ + 2e− → H 2
2 hydrogen atoms each gain 1 electron to yield 2 hydrogen atoms
with an oxidation state of 0.Notice that if we take the two half-reactions together, oxidation equals reduction:
In the oxidation half-reaction, iron loses 2 electrons. In the reduction half-
reaction, 2 hydrogen atoms each gain 1 electron; the total electron gain is 2.
That’s how these redox half-reactions work. Let’s try writing the half-reactions
for the following:
4NH 3 + 5O 2 → 4NO + 6H 2 O