Oxidation-Reduction Reactions, or redox reactions, are reactions in which
one reactant is oxidized and one reactant is reduced simultaneously. This
module demonstrates how to balance various redox equations.
Identifying Redox Reactions
The first step in balancing any redox reaction is determining whether or not it is
even an oxidation-reduction reaction. This requires that one and typically more
species changing oxidation states during the reaction. To maintain charge neutrality
in the sample, the redox reaction will entail both a reduction component and an
oxidation components. These are often separated into independent two
hypothetical half-reactions to aid in understanding the reaction. This requires
identifying which element is oxidized and which element is reduced. For example,
consider this reaction:
Cu(s)+2Ag+(aq)→Cu2+(aq)+2Ag(s)Cu(s)+2Ag+(aq)→Cu2+(aq)+2Ag(s)
The first step in determining whether the reaction is a redox reaction is to split the
equation into two hypothetical half-reactions. Let's start with the half-reaction
involving the copper atoms:
Cu(s)→Cu2+(aq)Cu(s)→Cu2+(aq)
The oxidation state of copper on the left side is 0 because it is an element on its
own. The oxidation state of copper on the right hand side of the equation is +2. The
copper in this half-reaction is oxidized as the oxidation states increases from 0
in CuCu to +2 in Cu2+Cu2+. Now consider the silver atoms
2Ag+(aq)→2Ag(s)2Ag+(aq)→2Ag(s)
In this half-reaction, the oxidation state of silver on the left side is a +1. The
oxidation state of silver on the right is 0 because it is an pure element. Because the
oxidation state of silver decreases from +1 to 0, this is the reduction half-reaction.
Consequently, this reaction is a redox reaction as both reduction and oxidation half-
reactions occur (via the transfer of electrons, that are not explicitly shown in
equations 2). Once confirmed, it often necessary to balance the reaction (the
reaction in equation 1 is balanced already though), which can be accomplished in
two ways because the reaction could take place in neutral, acidic or basic
conditions.
Balancing Redox Reactions
Balancing redox reactions is slightly more complex than balancing standard
reactions, but still follows a relatively simple set of rules. One major difference is the
necessity to know the half-reactions of the involved reactants; a half-reaction table
is very useful for this. Half-reactions are often useful in that two half reactions can
be added to get a total net equation. Although the half-reactions must be known to
complete a redox reaction, it is often possible to figure them out without having to
use a half-reaction table. This is demonstrated in the acidic and basic solution
examples. Besides the general rules for neutral conditions, additional rules must be
applied for aqueous reactions in acidic or basic conditions.
One method used to balance redox reactions is called the Half-Equation Method.
In this method, the equation is separated into two half-equations; one for oxidation
and one for reduction.
Half-Equation Method to Balance redox Reactions in AcidicAqueous Solutions
