17.1: Oxidation-reduction Reactions In a large and important class of reactions we find it useful to focus on the transfer of one or more electrons from one chemical moiety to another. Reactions in which electrons are transferred from one chemical moiety to another are called oxidation–reduction reactions, or redox reactions, for short. 17.2: Electrochemical Cells 17.3: Defining Oxidation States The definition of oxidation states predates our ability to estimate electron densities through quantum mechanical calculations. As it turns out, however, the ideas that led to the oxidation state formalism are directionally correct; atoms that have high positive oxidation states according to the formalism also have relatively high positive charges by quantum mechanical calculation. 17.4: Balancing Oxidation-reduction Reactions 17.5: Electrical Potential Electrical potential is measured in volts. If a system comprising one coulomb of charge passes through a potential difference of one volt, one joule of work is done on the system. Whether this represents an increase or a decrease in the energy of the system depends on the sign of the charge and on the sign of the potential difference. Electrical potential and gravitational potential are analogous. 17.6: Electrochemical Cells as Circuit Elements If the reaction between silver ions and copper metal is to occur, electrons must pass through the external circuit from the copper terminal to the silver terminal. An electron that is free to move in the presence of an electrical potential must move away from a region of more negative electrical potential and toward a region of more positive electrical potential. 17.7: The Direction of Electron Flow and its Implications The difference between electrolytic and galvanic cells lies in the direction of current flow and, correspondingly, the direction in which the cell reaction occurs. In a galvanic cell, a spontaneous chemical reaction occurs and this reaction determines the direction of current flow and the signs of the electrode potentials. In an electrolytic cell, the sign of the electrode potentials is determined by an applied potential source, which determines the direction of current flow. 17.8: Electrolysis and the Faraday Electrolytic cells are very important in the manufacture of products essential to our technology-intensive civilization. Only electrolytic processes can produce many materials, notably metals that are strong reducing agents. Aluminum and the alkali metals are conspicuous examples. Many manufacturing processes that are not themselves electrolytic utilize materials that are produced in electrolytic cells. These processes would not be possible if the electrolytic products were not available. 17.9: Electrochemistry and Conductivity From the considerations we have discussed, it is evident that any electrolytic cell involves a flow of electrons in an external circuit and a flow of ions within the materials comprising the cell. The function of the current collectors is to transfer electrons back and forth between the external circuit and the cell reagents. 17.10: The Standard Hydrogen Electrode (S.H.E) We also need to choose an arbitrary reference half-cell. The choice that has been adopted is the Standard Hydrogen Electrode, often abbreviated the S.H.E. The S.H.E. is defined as a piece of platinum metal, immersed in a unit-activity aqueous solution of a protonic acid, and over whose surface hydrogen gas, at unit fugacity, is passed continuously. These concentration choices make the electrode a standard electrode. 17.11: Half-reactions and Half-cells 17.12: Standard Electrode Potentials We adopt a very useful convention to tabulate the potential drops across standard electrochemical cells, in which one half-cell is the S.H.E. Since the potential of the S.H.E. is zero, we define the standard electrode potential, of any other standard half-cell (and its associated half-reaction) to be the potential difference when the half-cell operates spontaneously versus the S.H.E. The electrical potential of the standard half-cell determines both the magnitude and sign of the standard half-ce 17.13: Predicting the Direction of Spontaneous Change It is useful to associate the standard electrode potential with the half-reaction written as a reduction, that is, with the electrons written on the left side of the equation. We also establish the convention that reversing the direction of the half-reaction reverses the algebraic sign of its potential. When these conventions are followed, the overall reaction and the full-cell potential can be obtained by adding the corresponding half-cell information. 17.14: Cell Potentials and the Gibbs Free Energy 17.15: The Nernst Equation 17.16: The Nernst Equation for Half-cells 17.17: Combining two Half-cell Equations to Obtain a new Half-cell Equation 17.18: The Nernst Equation and the Criterion for Equilibrium 17.19: Problems Thumbnail: Schematic of Zn-Cu galvanic cell. (CC BY-SA 3.0; Ohiostandard ).​​