Galvanic Cells: Galvanic Cells

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 F₂, or H₂O₂) 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