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| title | chunk | source | category | tags | date_saved | instance |
|---|---|---|---|---|---|---|
| Electrolysis | 3/5 | https://en.wikipedia.org/wiki/Electrolysis | reference | science, encyclopedia | 2026-05-05T10:47:31.576904+00:00 | kb-cron |
=== Energy changes during electrolysis === The amount of electrical energy that must be added equals the change in Gibbs free energy of the reaction plus the losses in the system. The losses can (in theory) be arbitrarily close to zero, so the maximum thermodynamic efficiency equals the enthalpy change divided by the free energy change of the reaction. In most cases, the electric input is larger than the enthalpy change of the reaction, so some energy is released in the form of heat. In some cases, for instance, in the electrolysis of steam into hydrogen and oxygen at high temperature, the opposite is true and heat energy is absorbed. This heat is absorbed from the surroundings, and the heating value of the produced hydrogen is higher than the electric input.
=== Variations === Pulsating current results in products different from DC. For 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 == The chloralkali process is a large scale application of electrolysis. This technology supplies most of the chlorine and sodium hydroxide required by many industries. The cathode is a mixed metal oxide clad titanium anode (also called a dimensionally stable anode).
=== Electrofluorination === Many organofluorine compounds are produced by electrofluorination. One manifestation of this technology is the Simons process, which can be described as:
R3C–H + HF → R3C–F + H2 In the course of a typical synthesis, this reaction occurs once for each C–H bond in the precursor. The cell potential is maintained near 5–6 V. The anode, the electrocatalyst, is nickel-plated.
=== Hydrodimerization of acrylonitrile === Acrylonitrile is converted to adiponitrile on an industrial scale via electrocatalysis.
=== Electroplating and electrowinning processes === Purifying copper from refined copper. Electrometallurgy of aluminium, lithium, sodium, potassium, magnesium, calcium. 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) === In Electrochemical machining, 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.
=== Other === Production of sodium chlorate and potassium chlorate. Production of fuels such as hydrogen for spacecraft, nuclear submarines and vehicles. Rust removal and cleaning of old coins and other metallic objects.
== 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 SO2−4) 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 standard electrode potentials is shown below.
In terms of electrolysis, this table should be interpreted as follows:
Moving down the table, E° becomes more positive, and species on the left are more likely to be reduced: for example, zinc ions are more likely to be reduced to zinc metal than sodium ions are to be reduced to sodium metal. Moving up the table, E° becomes more negative, and species on the right are more likely to be oxidized: for example, sodium metal is more likely to be oxidized to sodium ions than zinc metal is to be oxidized to zinc ions. Using the Nernst equation the electrode potential can be calculated for a specific concentration of ions, temperature and the number of electrons involved. For pure water (pH 7):