CHEMISTRY THE CENTRAL SCIENCE
CHAPTER SUMMARY AND KEY TERMS
INTRODUCTION AND SECTION 20.1 In this chapter we have focused on electrochemistry, the branch of chemistry that relates electricity and chemical reactions. Electrochemistry involves oxidation-reduction reactions, also called redox reactions. These reactions involve a change in the oxidation state of one or more elements. In every oxidation-reduction reaction one substance is oxidized (its oxidation state, or number, increases) and one substance is reduced (its oxidation state, or number, decreases). The substance that is oxidized is referred to as a reducing agent, orreductant, because it causes the reduction of some other substance. Similarly, the substance that is reduced is referred to as an oxidizing agent, or oxidant, because it causes the oxidation of some other substance.
SECTION 20.2 An oxidization-reduction reaction can be balanced by dividing the reaction into two half-reactions, one for oxidation and one for reduction. A half-reaction is a balanced chemical equation that includes electrons. In oxidation half-reactions the electrons are on the product (right) side of the equation; we can envision that these electrons are transferred from a substance when it is oxidized. In reduction half-reactions the electrons are on the reactant (left) side of the equation. Each half-reaction is balanced separately, and the two are brought together with proper coefficients to balance the electrons on each side of the equation, so the electrons cancel when the half-reactions are added.
SECTION 20.3 A voltaic (or galvanic) cell uses a spontaneous oxidation-reduction reaction to generate electricity. In a voltaic cell the oxidation and reduction half-reactions often occur in separate half-cells. Each half-cell has a solid surface called an electrode, where the half-reaction occurs. The electrode where oxidation occurs is called the anode; reduction occurs at the cathode. The electrons released at the anode flow through an external circuit (where they do electrical work) to the cathode. Electrical neutrality in the solution is maintained by the migration of ions between the two half-cells through a device such as a salt bridge.
SECTION 20.4 A voltaic cell generates an electromotive force (emf) that moves the electrons from the anode to the cathode through the external circuit. The origin of emf is a difference in the electrical potential energy of the two electrodes in the cell. The emf of a cell is called its cell potential, Ecell, and is measured in volts (1 V = 1 J/C). The cell potential under standard conditions is called the standard emf, or the standard cell potential, and is denoted .
A standard reduction potential, , can be assigned for an individual half-reaction. This is achieved by comparing the potential of the half-reaction to that of the standard hydrogen electrode (SHE), which is defined to have = 0 V and is based on the following half-reaction:
The standard cell potential of a voltaic cell is the difference between the standard reduction potentials of the half-reactions that occur at the cathode and the anode: = (cathode) – (anode). The value of is positive for a voltaic cell.
For a reduction half-reaction, is a measure of the tendency of the reduction to occur; the more positive the value for , the greater the tendency of the substance to be reduced. Thus, provides a measure of the oxidizing strength of a substance. Substances that are strong oxidizing agents produce products that are weak reducing agents, and vice versa.
SECTION 20.5 The emf, E, is related to the change in the Gibbs free energy, ΔG: ΔG = –nFE, where n is the number of electrons transferred during the redox process and F is Faraday's constant, defined as the quantity of electrical charge on one mole of electrons: F = 96,485 C/mol. Because E is related to ΔG, the sign of E indicates whether a redox process is spontaneous: E > 0 indicates a spontaneous process, and E < 0 indicates a nonspontaneous one. Because ΔG is also related to the equilibrium constant for a reaction (ΔG = –RT ln K), we can relate E to K as well.
The maximum amount of electrical work produced by a voltaic cell is given by the product of the total charge delivered, nF, and the emf, E: wmax = –nFE. The watt is a unit of power: 1 W = 1 J/s. Electrical work is often measured in kilowatt-hours.
SECTION 20.6 The emf of a redox reaction varies with temperature and with the concentrations of reactants and products. The Nernst equation relates the emf under nonstandard conditions to the standard emf and the reaction quotient Q:
The factor 0.0592 is valid when T = 298K. A concentration cell is a voltaic cell in which the same half-reaction occurs at both the anode and cathode but with different concentrations of reactants in each half-cell. At equilibrium, Q = K and E = 0.
SECTION 20.7 A battery is a self-contained electrochemical power source that contains one or more voltaic cells. Batteries are based on a variety of different redox reactions. Several common batteries were discussed. The lead-acid battery, the nickel-cadmium battery, the nickel-metal-hydride battery, and the lithium-ion battery are examples of rechargeable batteries. The common alkaline dry cell is not rechargeable. Fuel cells are voltaic cells that utilize redox reactions in which reactants such as H2 have to be continuously supplied to the cell to generate voltage.
SECTION 20.8 Electrochemical principles help us understand corrosion, undesirable redox reactions in which a metal is attacked by some substance in its environment. The corrosion of iron into rust is caused by the presence of water and oxygen, and it is accelerated by the presence of electrolytes, such as road salt. The protection of a metal by putting it in contact with another metal that more readily undergoes oxidation is called cathodic protection. Galvanized iron, for example, is coated with a thin layer of zinc; because zinc is oxidized more readily than iron, the zinc serves as a sacrificial anode in the redox reaction.
SECTION 20.9 An electrolysis reaction, which is carried out in an electrolytic cell, employs an external source of electricity to drive a nonspontaneous electrochemical reaction. The current-carrying medium within an electrolytic cell may be either a molten salt or an electrolyte solution. The products of electrolysis can generally be predicted by comparing the reduction potentials associated with possible oxidation and reduction processes. The electrodes in an electrolytic cell can be active, meaning that the electrode can be involved in the electrolysis reaction. Active electrodes are important in electroplating and in metallurgical processes.
The quantity of substances formed during electrolysis can be calculated by considering the number of electrons involved in the redox reaction and the amount of electrical charge that passes into the cell. The amount of electrical charge is measured in coulombs and is related to the magnitude of the current and the time it flows (1 C = 1 A-s). Electrometallurgy is the use of electrolytic methods to prepare or purify a metallic element. Aluminum is obtained in the Hall-Héroult process by electrolysis of Al2O3 in molten cryolite (Na3AlF6).
• Identify oxidation, reduction, oxidizing agent, and reducing agent in a chemical equation. (Section 20.1)
• Complete and balance redox equations using the method of half-reactions. (Section 20.2)
• Sketch a voltaic cell and identify its cathode, anode, and the directions that electrons and ions move. (Section 20.3)
• Calculate standard emfs (cell potentials), , from standard reduction potentials. (Section 20.4)
• Use reduction potentials to predict whether a redox reaction is spontaneous. (Section 20.4)
• Relate to ΔG° and equilibrium constants. (Section 20.5)
• Calculate emf under nonstandard conditions. (Section 20.6)
• Describe the components of common batteries and fuel cells. (Section 20.7)
• Explain how corrosion occurs and how it is prevented by cathodic protection. (Section 20.8)
• Describe the reactions in electrolytic cells. (Section 20.9)
• Relate amounts of products and reactants in redox reactions to electrical charge. (Section 20.9)
Relating standard emf to standard reduction potentials of the reduction (cathode) and oxidation (anode) half-reactions
Relating free-energy change and emf
The Nernst equation, expressing the effect of concentration on cell potential