Lecture 19 summary

In this lecture, while we continue learning the principles of electrochemistry, we begin to turn our attention more toward applications. The key relationship between the change in standard reduction potential, DE°, and the equilibrium constant is derived from the Nernst equation, and the result applied to examples where we construct a complete redox reaction by combining two half reactions.

Procedure for writing a net redox reaction

To get a complete, balanced redox reaction, and obtain its DE° value, written in the direction favoring products at equilibrium (DE° > 0), we combine half reactions using the following "recipe::

Step 1. Write the two reduction half-reactions along with their associated E° values (Appendix C). Multiply the half-reactions as necessary so that they each contain the same number of electrons. When you multiply a reaction by any number, do not multiply E°.

Step 2. To write a balanced net cell reaction, subtract the half-reaction with the lower E° value (E°_L) from the other half-reaction. (This operation is equivalent to reversing the lower E° value half-reaction and adding.)

Step 3. Find the net change in standard reduction potential DE° by subtraction: DE° = E°_H – E°_L.

Compare these rules with those given in Harris on p.298. Here we are not necessarily concerned about relating our redox reaction to a galvanic cell, but are mainly concerned about working with half-reactions, combining them in various ways, calculating their DE° values. We will deal explicitly with applying the Nernst equation for a complete redox reaction below, and as we will see, the DE° value we derive for a redox reaction is used to calculate its equilibrium constant.