Redox Potentials and Galvanic Cells

Abstract

In the last chapter, we learned a lot about Galvani potential differences across different individual interfaces and the usefulness of these potential differences, but we did not get to know how they can be measured. The problem is that it is impossible to measure the Galvani potential difference across a single interface in a half-cell directly, because an electrolyte’s contact to the conductors of an electric measuring device requires a second electrode and this produces a new interface with an additional Galvani potential difference. The way out of the dilemma is the use of always the same reference half-cell, the standard hydrogen electrode, so that the measured voltage of the galvanic cell is only determined by the measuring half-cell. In this way, we obtain the so-called redox potential of a half-cell which represents, just like the electron potential, a measure of the strengths of reducing or oxidizing agents. The redox potentials under standard conditions are often compiled according to their values in a table, the electrochemical series. The combination of two arbitrary half-cells results in a galvanic cell. The reversible cell voltage of such a cell, meaning the cell voltage in equilibrium, can be described by Nernst’s equation and used to predict the chemical drive, the equilibrium constant, and other thermodynamic properties of chemical reactions. Subsequently, some technically important galvanic cells will be discussed, which yield usable energy due to the spontaneous chemical reactions running inside them. In closing, the technique of potentiometry and the corresponding potentiometric titration is presented. This electroanalytical method uses the concentration dependence of the reversible cell voltage for quantitative analysis of ions.