example, pulsing increases the ratio of ozone to oxygen produced at the anode in the electrolysis of an aqueous acidic solution such as dilute sulphuric acid. Electrolysis of ethanol with pulsed current evolves an aldehyde instead of primarily an acid. Related processes Galvanic cells and batteries use spontaneous, energy-releasing redox reactions to generate an electrical potential that provides useful power. When a secondary battery is charged, its redox reaction is run in reverse and the system can be considered as an electrolytic cell. Industrial uses Electrometallurgy of aluminium, lithium, sodium, potassium, magnesium, calcium, and in some cases copper. Production of chlorine and sodium hydroxide, called the Chloralkali process. Production of sodium chlorate and potassium chlorate. Production of perfluorinated organic compounds such as trifluoroacetic acid by the process of electrofluorination. Purifying copper from refined copper. Production of fuels such as hydrogen for spacecraft, nuclear submarines and vehicles. Rust removal and cleaning of old coins and other metallic objects. Manufacturing processes In manufacturing, electrolysis can be used for: Electroplating, where a thin film of metal is deposited over a substrate material. Electroplating is used in many industries for either functional or decorative purposes, as in-vehicle bodies and nickel coins. Electrochemical machining (ECM), where an electrolytic cathode is used as a shaped tool for removing material by anodic oxidation from a workpiece. ECM is often used as a technique for deburring or for etching metal surfaces like tools or knives with a permanent mark or logo. Competing half-reactions in solution electrolysis Using a cell containing inert platinum electrodes, electrolysis of aqueous solutions of some salts leads to the reduction of the cations (such as metal deposition with, for example, zinc salts) and oxidation of the anions (such as the evolution of bromine with bromides). However, with salts of some metals (such as sodium) hydrogen is evolved at the cathode, and for salts containing some anions (such as sulfate ) oxygen is evolved at the anode. In both cases, this is due to water being reduced to form hydrogen or oxidized to form oxygen. In principle, the voltage required to electrolyze a salt solution can be derived from the standard electrode potential for the reactions at the anode and cathode. The standard electrode potential is directly related to the Gibbs free energy, ΔG, for the reactions at each electrode and refers to an electrode with no current flowing. An extract from the table of