Electrochemistry

History

Investigations into electrical phenomena began in the sixteenth century. The English scientist William Gilbert spent 17 years experimenting with magnetism and electricity, becoming known as the Father of Magnetism for his work on magnets. In 1663, the German physicist Otto von Guericke built the first electric generator, which produced static electricity by friction; it consisted of a rotating sulfur ball cast inside a glass globe. By the mid-18th century, the French chemist Charles François de Cisternay du Fay had identified two types of static electricity, proposing a two-fluid theory with vitreous (positive) and resinous (negative) electricity, later opposed by Benjamin Franklin's one-fluid theory.

In 1785, Charles-Augustin de Coulomb formulated the law of electrostatic attraction. The late 18th century saw the Italian physician Luigi Galvani mark the birth of electrochemistry with his 1791 essay on the effect of electricity on muscular motion, in which he proposed a nerveo-electrical substance and concluded that animal tissue contained an innate vital force he called animal electricity, distinct from natural and artificial forms.

Principles

The term redox stands for reduction-oxidation, referring to electrochemical processes involving electron transfer that changes the oxidation state of a molecule or ion. Oxidation is the loss of electrons, and reduction is the gain of electrons; these two processes always occur together. The oxidation state is the hypothetical charge an atom would have if all bonds to different elements were 100% ionic. An atom or ion that gives up an electron has its oxidation state increase, while the recipient's oxidation state decreases.

For example, sodium donates one electron to chlorine, attaining an oxidation state of +1, while chlorine accepts the electron and its oxidation state becomes -1. The substance that loses electrons is the reducing agent, and the one that accepts electrons is the oxidizing agent. Despite the name, an oxidation reaction does not necessarily involve oxygen; fluorine, for instance, is an even stronger oxidant than oxygen. Mnemonic devices such as OIL RIG (Oxidation Is Loss, Reduction Is Gain) and LEO the lion says GER (Lose Electrons: Oxidation, Gain Electrons: Reduction) help remember these definitions.

Corrosion

Corrosion is an electrochemical process that manifests as rust on iron or tarnish on copper and their alloys. For iron to rust, the metal must be in contact with oxygen and water. One area of the metal surface acts as the anode, where iron gives up electrons: Fe(s) → Fe2+(aq) + 2 e−. These electrons reduce oxygen in the atmosphere to water at the cathode: O2(g) + 4 H+(aq) + 4 e− → 2 H2O(l).

The overall reaction is 2 Fe(s) + O2(g) + 4 H+(aq) → 2 Fe2+(aq) + 2 H2O(l), with a standard emf of 1.67 V. Iron corrosion occurs in an acid medium; H+ ions come from carbonic acid formed by carbon dioxide and water. The Fe2+ ions further oxidize to form iron(III) oxide hydrate, or rust, represented as Fe2O3·xH2O.

An electrolyte facilitates oxidation, explaining why rusting is quicker in salt water. Copper develops a green-blue patina of copper carbonate, while silver forms black silver sulfide when exposed to sulfur. Gold and platinum resist oxidation under normal conditions.

Electrolysis

Electrolysis uses an external source of electrical energy to drive a non-spontaneous chemical reaction in an electrolytic cell. Molten sodium chloride can be electrolyzed in a Downs cell to produce metallic sodium and chlorine gas. At the anode, chloride ions are oxidized: 2 Cl−(l) → Cl2(g) + 2 e−; at the cathode, sodium ions are reduced: 2 Na+(l) + 2 e− → 2 Na(l). The overall reaction is 2 Na+(l) + 2 Cl−(l) → 2 Na(l) + Cl2(g), with an emf of approximately −4 V, requiring at least 4 V from the power supply.

Water can also be decomposed into hydrogen and oxygen by applying an external voltage in an electrolytic cell with inert platinum electrodes. The overall reaction is 2 H2O(l) → 2 H2(g) + O2(g). This process does not occur spontaneously because the Gibbs free energy change is very positive (about 474.4 kJ). An electrolyte such as sodium chloride or sulfuric acid is needed for the reaction to proceed at practical voltages.

Did You Know?

• An electrochemical cell produces an electric current from energy released by a spontaneous redox reaction and includes the Galvanic cell or Voltaic cell.
• The electrical energy from an electrochemical cell can be calculated as the product of the cell's emf and the electric charge transferred, where charge equals the number of electrons transferred times Faraday's constant.