In chemistry and manufacturing, electrolysis is a technique that uses direct electric current (DC) to drive an otherwise non-spontaneous biological and physical reaction. Electrolysis is commercially important as a stage in the separation of elements from naturally occurring sources such as ores using an electrolytic cell. The voltage that is needed for electrolysis to occur is called the decomposition potential. The word "lysis" means to separate or break, so in terms, electrolysis would mean "breakdown via electricity."
Etymology The word "electrolysis" was introduced by Michael Faraday in 1834, using the Greek words "amber", which since the 17th century was associated with electrical phenomena, and meaning "dissolution". Nevertheless, electrolysis, as a tool to study chemical reactions and obtain pure elements, precedes the coinage of the term and formal description by Faraday.
History In the early nineteenth century, William Nicholson and Anthony Carlisle sought to further Volta's experiments. They attached two wires to either side of a voltaic pile and placed the other ends in a tube filled with water. They noticed that when the wires were brought together, each wire produced bubbles; one type produced hydrogen bubbles, and the other produced oxygen bubbles.
In 1785 a Dutch scientist named Martin van Marum created an electrostatic generator that he used to reduce tin, zinc and antimony from their salts using a process later known as electrolysis. Though he unknowingly produced electrolysis, it was not until 1800 that William Nicholson and Anthony Carlisle discovered how electrolysis works.
In 1791, Luigi Galvani experimented with frog legs. He claimed that placing animal muscle between two dissimilar metal sheets resulted in electricity. Responding to these claims, Alessandro Volta conducted his own tests. This would give insight to Humphry Davy's ideas on electrolysis. During preliminary experiments, Davy hypothesized that when two elements combine to form a compound, electrical energy is released. Davy would go on to create Decomposition Tables from his preliminary experiments on Electrolysis. The Decomposition Tables would give insight on the energies needed to break apart certain compounds.
In 1817, Johan August Arfwedson determined there was another element, lithium, in some of his samples; however, he could not isolate the component. In 1821, William Thomas Brande used electrolysis to single it out, and two years later, he streamlined the process using lithium chloride and potassium chloride with electrolysis to produce lithium and lithium hydroxide.
During the later years of Davy's research, Faraday became his assistant. While studying the process of electrolysis under Davy, Faraday discovered two laws of electrolysis.
During the time of Maxwell and Faraday, concerns came about for electropositive and electronegative activities.
In November 1875, Paul Émile Lecoq de Boisbaudran discovered gallium using electrolysis of gallium hydroxide, producing 3.4 mg of gallium. The following December, he presented his discovery of gallium to the Académie des sciences in Paris.
On June 26, 1886, Ferdinand Frederick Henri Moissan finally felt comfortable performing electrolysis on anhydrous hydrogen fluoride to create a gaseous fluorine pure element. Before he used hydrogen fluoride, Moissan used fluoride salts with electrolysis. On June 28, 1886, he performed his experiment in front of the Académie des sciences to show his discovery of the new element fluorine. While trying to find elemental fluorine through electrolysis of fluoride salts, many chemists perished, including Paulin Louyet and Jérôme Nicklès.
In 1886 Charles Martin Hall from America and Paul Héroult from France both filed for American patents for the electrolysis of aluminum, with Héroult submitting his in May, and Hall in July. Hall was able to get his patent by proving through letters to his brother and family evidence that his method was discovered before the French patent was submitted. This became known as the Hall–Héroult process, which benefited many industries because aluminum's price then dropped from four dollars to thirty cents per pound.
In 1902 Polish engineer and inventor Stanisław Łaszczyński filed for and obtained Polish patent for the electrolysis of copper and zinc.
Timeline 1785 – Martinus van Marum's electrostatic generator was used to reduce tin, zinc, and antimony from their salts using electrolysis.The Supplement (1803 edition) to Encyclopædia Britannica 3rd edition (1797), volume 1, page 225, "Mister Van Marum, by means of his great electrical machine, decomposed the calces of tin, zinc, and antimony, and resolved them into their respective metals and oxygen" and gives as a reference Journal de Physiques, 1785. 1789 – Adriaan Paets van Troostwijk and Jan Rudolph Deiman conducted early experiments on water electrolysis using an electrostatic generator, producing small amounts of hydrogen and oxygen. 1800 – William Nicholson and Anthony Carlisle (and also Johann Ritter), demonstrated the first sustained electrolysis of water into hydrogen and oxygen using a voltaic pile. 1808 – Potassium (1807), sodium (1807), barium, calcium and magnesium were discovered by Humphry Davy using electrolysis. 1821 – Lithium was discovered by the English chemist William Thomas Brande, who obtained it by electrolysis of lithium oxide. 1834 – Michael Faraday published his two laws of electrolysis, provided a mathematical explanation for them, and introduced terminology such as electrode, electrolyte, anode, cathode, anion, and cation. 1875 – Paul Émile Lecoq de Boisbaudran discovered gallium using electrolysis. 1886 – Fluorine was discovered by Henri Moissan using electrolysis. 1886 – Hall–Héroult process developed for making aluminium. 1890 – Castner–Kellner process developed for making sodium hydroxide. 1902 – Stanisław Łaszczyński obtained copper using electrolysis. 1930 – Development of the modern chlor-alkali process (electrolysis of brine to produce chlorine and sodium hydroxide), which became an important industrial method.Electrochemical Society. (n.d.). Metals & Chemicals. Retrieved February 2, 2025, fromElectrochemical Society. (n.d.). Metals & Chemicals. Retrieved https://www1.eere.energy.gov/manufacturing/resources/chemicals/pdfs/profile_chap6.pdf
Overview Electrolysis is the passing of a direct electric current through an electrolyte which is producing chemical reactions at the electrodes and decomposition of the materials.
The main components required to achieve electrolysis are an electrolyte, electrodes, and an external power source. A partition (e.g. an ion-exchange membrane or a salt bridge) is optional to keep the products from diffusing to the vicinity of the opposite electrode.
The electrolyte is a chemical substance which contains free ions and carries electric current (e.g. an ion-conducting polymer, solution, or an ionic liquid compound). If the ions are not mobile, as in most solid salts, then electrolysis cannot occur. A liquid electrolyte is produced by: Solvation or reaction of an ionic compound with a solvent (such as water) to produce mobile ions An ionic compound melted by heating
The electrodes are immersed separated by a distance such that a current flows between them through the electrolyte and are connected to the power source which completes the electrical circuit. A direct current supplied by the power source drives the reaction causing ions in the electrolyte to be attracted toward the respective oppositely charged electrode.
Electrodes of metal, graphite and semiconductor material are widely used. Choice of suitable electrode depends on chemical reactivity between the electrode and electrolyte and manufacturing cost. Historically, when non-reactive anodes were desired for electrolysis, graphite (called plumbago in Faraday's time) or platinum were chosen. They were found to be some of the least reactive materials for anodes. Platinum erodes very slowly compared to other materials, and graphite crumbles and can produce carbon dioxide in aqueous solutions but otherwise does not participate in the reaction. Cathodes may be made of the same material, or they may be made from a more reactive material, since anode wear is greater due to oxidation at the anode.
Process of electrolysis The key process of electrolysis is the interchange of atoms and ions by the removal or addition of electrons due to the applied potential. The desired products of electrolysis are often in a different physical state from the electrolyte and can be removed by mechanical processes (e.g. by collecting gas above an electrode or precipitating a product out of the electrolyte).
The quantity of the products is proportional to the current, and when two or more electrolytic cells are connected in series to the same power source, the products produced in the cells are proportional to their equivalent weight. These are known as Faraday's laws of electrolysis.
Each electrode attracts ions that are of the opposite charge. Positively charged ions (cations) move towards the electron-providing (negative) cathode. Negatively charged ions (anions) move towards the electron-extracting (positive) anode. In this process, electrons are effectively introduced at the cathode as a reactant and removed at the anode as a product. In chemistry, the loss of electrons is called oxidation, while electron gain is called reduction.
When neutral atoms or molecules, such as those on the surface of an electrode, gain or lose electrons they become ions and may dissolve in the electrolyte and react with other ions.
When ions gain or lose electrons and become neutral, they will form compounds that separate from the electrolyte. Positive metal ions like copper (ii) (Cu2+) deposit onto the cathode in a layer. The terms for this are electroplating, electrowinning, and electrorefining.
When an ion gains or loses electrons without becoming neutral, its electronic charge is altered in the process.
For example, the electrolysis of brine produces hydrogen and chlorine gases which bubble from the electrolyte and are collected. The initial overall reaction is thus:
2 NaCl + 2 H2O → 2 NaOH + H2 + Cl2
The reaction at the anode results in chlorine gas from chlorine ions: 2 Cl− → Cl2 + 2 e−
The reaction at the cathode results in hydrogen gas and hydroxide ions: 2 H2O + 2 e− → H2 + 2 OH−
Without a partition between the electrodes, the OH− ions produced at the cathode are free to diffuse throughout the electrolyte to the anode. As the electrolyte becomes more basic due to the production of OH−, less Cl2 emerges from the solution as it begins to react with the hydroxide, producing hypochlorite (ClO−) at the anode:
Cl2 + 2 NaOH → NaCl + NaClO + H2O
The more opportunity the Cl2 has to interact with NaOH in the solution, the less Cl2 emerges at the surface of the solution and the faster the production of hypochlorite progresses. This depends on factors such as solution temperature, the amount of time the Cl2 molecule is in contact with the solution, and concentration of NaOH.
Likewise, as hypochlorite increases in concentration, chlorates are produced from them:
3 NaClO → NaClO3 + 2 NaCl
Other reactions occur, such as the self-ionization of water and the decomposition of hypochlorite at the cathode, the rate of the latter depends on factors such as diffusion and the surface area of the cathode in contact with the electrolyte.
Decomposition potential Decomposition potential or decomposition voltage refers to the minimum voltage (difference in electrode potential) between anode and cathode of an electrolytic cell that is needed for electrolysis to occur.