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.
Figure %: Determining the Oxidation States 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:
Atoms in elemental form have oxidation states of zero.
The charge on a monoatomic ion is equivalent to its charge.
Hydrogen has an oxidation state of +1 when bonded to non-metals and -1
when bonded
to metals.
F, because it only forms one bond and is the most electronegative
element, has an
oxidation state of -1.
O, unless bonded to F or itself, has an oxidation state of -2.
The sum of all oxidation states in a compound must equal the total
charge on the species.
For hydrogen peroxide, we have seen that both O's have -1 oxidation states
and both H's
have +1 oxidation states, the sum of which is zero--the charge on hydrogen
peroxide.
Balancing Redox Reactions
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.
