In order to keep track of the electron flow in redox reactions we will define the concept of oxidation state and provide some rules for determining the oxidation states of atoms within a compound. The oxidation state of an atom in a covalent compound is an imaginary charge assigned to that atom if all the electrons in its bonds were completely given to the more electronegative atom in the bond. Of course, if two atoms of the same element are bonded together, then the two electrons in its bond are shared equally. In ionic compounds, the charge on the ion is equal to its oxidation number. shows, using hydrogen peroxide as an example, how one calculates the oxidation state of an atom in a covalent compound.
The above procedure for assigning oxidation states leads to the following useful observations about oxidation states which you should verify given the above discussion:
Oxidation-reduction reactions, also called redox reactions, involve the transfer of electrons from one species to another. That electron transfer causes a change in oxidation state for both reactive partners. The reducing agent is oxidized, meaning that its oxidation number increases due to the loss of one or more electrons. The oxidizing agent is reduced, meaning that its oxidation number has decreased due to the gain of one or more electrons. For example, the between permanganate and iron metal in acidic aqueous solution involves the transfer of five electrons to each permanganate (the oxidizing agent) from iron (the reducing agent).
Balancing redox equations by inspection is quite difficult, as you must take into account not only the mass balance but also the charge balance in the equation. To aid in this task a set of rules, called the Half-Reaction Method, has been devised. The following rules work for reactions performed in acidic or in basic solution.
1. Separate oxidation and reduction half-reactions:
2. Balance all atoms except for hydrogen and oxygen in each half-reaction.
In this example
they are already balanced.