Because the SHE has a potential of exactly zero volts, as defined above,
the reaction:

has a value of 0.34 V for its E^{o} (recall that
E^{o}_{cell}
= E^{o}_{SHE} + E^{o}). Fortunately, every important
reduction potential has been measured and tabulated. Useful lists of
reduction potentials
are available in most introductory chemistry texts, including yours. In
this SparkNote, all
potentials will be given to you if you need them.

Those tables of standard reduction potentials list all half-reactions as
reductions. Half-
reactions with the largest reduction potential are placed at the top of the
list and the smallest
(most negative) reduction potentials are at the bottom. Those species on
the left-hand side
of the equations at the top of the list are the most easily reduced (like
F_{2}, or
H_{2}O_{2}) and those at the bottom are the least readily
reduced (like
Li^{+}).

Take a look at the list of standard reduction potentials in your chemistry
text. An intuitive
trend should be obvious when looking at the data--electronegative species
(those with the
greatest attraction for electrons) are easily reduced, i.e. given an
electron. The most
electronegative atom, F, has the largest reduction potential whereas one of
the least
electronegative atoms, Li, has the smallest reduction potential.

Adding Standard Reduction Potentials

By compiling a list of standard reduction potentials of all possible
reductions, one can, at
least in theory, calculate the cell potential,
E^{o}_{cell}, of any
arbitrary redox reaction. By knowing the sign of
E^{o}_{cell}, we
can predict whether a reaction is spontaneous at standard conditions. If
E^{o}_{cell} is positive, then the reaction is spontaneous.
Conversely, if E^{o}_{cell} is negative, then the reaction
is non-spontaneous as written but spontaneous in the reverse direction (see
Thermodynamics, Electrical Work and Cell Potential for an
explanation of why that is so).

To compute the cell potential of a reaction at standard conditions,
E^{o}_{cell}, you do not need to balance the equation of
your redox
reaction. However, as we will learn in Thermodynamics of
Electrochemistry, if the reaction is not conducted at
standard state, then
it is essential to balance the redox reaction to compute its cell
potential. For now, let's
assume that all reactions are conducted at standard conditions unless
otherwise specified.

When asked to compute the cell potential for a reaction, you will need to
be able to separate
the overall reaction into its oxidation and reduction half-reactions as in
.

Once those half reactions are separated, then find the reduction potential
for the reaction
written as a reduction. As you can see in , one
reaction is written
as an oxidation. For that reaction, you need to calculate its oxidation
potential. To
calculate an oxidation potential, simply reverse the sign of the
E^{o} for the
corresponding reduction reaction (just the oxidation written in the
opposite direction).
Simply add the reduction potential of the reduction and the oxidation
potential of the
oxidation to calculate the E^{o}_{cell}. It is important
to note here that
E^{o}'s are intrinsic properties of reactions and, therefore, do
not depend on the
stoichiometry of the reaction. That means that you DO NOT multiply the
E^{o} of a reaction by the coefficient used to balance the overall
redox reaction.
A proof of that point is provided in Thermodynamics of
Electrochemistry#. Multiplying
the value of E^{o} for a
half-reaction is the number one mistake made in calculating
E^{o}_{cell}. Please, don't let that happen to you!
Simply read off the values of E^{o} for the oxidation and reduction half-reactions
and add those two values together, as in .
/PARAGRAPH