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Illustration of a Hofmann electrolysis apparatus used in a school laboratory Electrolysis Apparatus.png
Illustration of a Hofmann electrolysis apparatus used in a school laboratory

In chemistry and manufacturing, electrolysis is a technique that uses direct electric current (DC) to drive an otherwise non-spontaneous chemical 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".



The word "electrolysis" was introduced by Michael Faraday in 1834, [1] using the Greek words ἤλεκτρον [ɛ̌ːlektron] "amber", which since the 17th century was associated with electrical phenomena, and λύσις [lýsis] 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.


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 when the wires were brought together that each wire produced bubbles. One type was hydrogen, the other was oxygen. [2]

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 when William Nicholson and Anthony Carlisle discovered how electrolysis works. [3]

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. [4] [5] This would give insight to Humphry Davy's ideas on electrolysis. During preliminary experiments, Humphry Davy hypothesized that when two elements combine together to form a compound, electrical energy is released. Humphry 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. [6]

In 1817 Johan August Arfwedson determined there was another element, lithium, in some of his samples; however, he could not isolate the component. It was not until 1821 when William Thomas Brande used electrolysis to single it out. Two years later, he streamlined the process using lithium chloride and potassium chloride with electrolysis to produce lithium and lithium hydroxide. [7] [8]

During the later years of Humphry Davy's research, Michael Faraday became his assistant. While studying the process of electrolysis under Humphry Davy, Michael Faraday discovered two laws of electrolysis. [5]

During the time of Maxwell and Faraday, concerns came about[ vague ] for electropositive and electronegative activities. [9]

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. [10]

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, Henri Moissan used fluoride salts with electrolysis. Thus on June 28, 1886, he performed his experiment in front of the Académie des sciences to show his discovery of the new element fluorine. [11] While trying to find elemental fluorine through electrolysis of fluoride salts, many chemists perished including Paulin Louyet and Jérôme Nicklès. [12]

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. [13] 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. [14] 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. [15]



Electrolysis is the passing of a direct electric current through an electrolyte 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 a ionic liquid compound). If the ions are not mobile, as in most solid salts, then electrolysis cannot occur. A liquid electrolyte is produced by:

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. [18] 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 one 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 current. 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 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: [19]

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. [20]

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. [21]

The voltage at which electrolysis is thermodynamically preferred is the difference of the electrode potentials as calculated using the Nernst equation. Applying additional voltage, referred to as overpotential, can increase the rate of reaction and is often needed above the thermodynamic value. It is especially necessary for electrolysis reactions involving gases, such as oxygen, hydrogen or chlorine.

Oxidation and reduction at the electrodes

Oxidation of ions or neutral molecules occurs at the anode. For example, it is possible to oxidize ferrous ions to ferric ions at the anode:

(aq) → Fe3+
(aq) + e

Reduction of ions or neutral molecules occurs at the cathode. It is possible to reduce ferricyanide ions to ferrocyanide ions at the cathode:

+ e → Fe(CN)4-

Neutral molecules can also react at either of the electrodes. For example: p-benzoquinone can be reduced to hydroquinone at the cathode:

P-Benzochinon.svg + 2 e + 2 H+ Hydrochinon2.svg

In the last example, H+ ions (hydrogen ions) also take part in the reaction and are provided by the acid in the solution, or by the solvent itself (water, methanol, etc.). Electrolysis reactions involving H+ ions are fairly common in acidic solutions. In aqueous alkaline solutions, reactions involving OH (hydroxide ions) are common.

Sometimes the solvents themselves (usually water) are oxidized or reduced at the electrodes. It is even possible to have electrolysis involving gases, e.g. by using a gas diffusion electrode.

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.


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. [22] Electrolysis of ethanol with pulsed current evolves an aldehyde instead of primarily an acid. [23]

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

Hall-Heroult process for producing aluminium Hall-heroult-kk-2008-12-31.png
Hall-Heroult process for producing aluminium

Manufacturing processes

In manufacturing, electrolysis can be used for:

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−
) 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.

Half-reaction E° (V)Ref.
Na + + e Na(s)−2.71 [24]
Zn 2+ + 2 e Zn(s)−0.7618 [25]
2 H+ + 2 e H2(g)≡ 0 [25]
Br2(aq) + 2 e 2 Br+1.0873 [25]
O2(g) + 4 H+ + 4 e 2 H2O+1.23 [24]
Cl2(g) + 2 e 2 Cl+1.36 [24]
+ 2 e 2 SO2−
+2.07 [24]

In terms of electrolysis, this table should be interpreted as follows:

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):

Comparable figures calculated in a similar way, for 1 M zinc bromide, ZnBr2, are −0.76 V for the reduction to Zn metal and +1.10 V for the oxidation producing bromine. The conclusion from these figures is that hydrogen should be produced at the cathode and oxygen at the anode from the electrolysis of water—which is at variance with the experimental observation that zinc metal is deposited and bromine is produced. [26] The explanation is that these calculated potentials only indicate the thermodynamically preferred reaction. In practice, many other factors have to be taken into account such as the kinetics of some of the reaction steps involved. These factors together mean that a higher potential is required for the reduction and oxidation of water than predicted, and these are termed overpotentials. Experimentally it is known that overpotentials depend on the design of the cell and the nature of the electrodes.

For the electrolysis of a neutral (pH 7) sodium chloride solution, the reduction of sodium ion is thermodynamically very difficult and water is reduced evolving hydrogen leaving hydroxide ions in solution. At the anode the oxidation of chlorine is observed rather than the oxidation of water since the overpotential for the oxidation of chloride to chlorine is lower than the overpotential for the oxidation of water to oxygen. The hydroxide ions and dissolved chlorine gas react further to form hypochlorous acid. The aqueous solutions resulting from this process is called electrolyzed water and is used as a disinfectant and cleaning agent.

Electrolysis of carbon dioxide

The electrochemical reduction or electrocatalytic conversion of CO2 can produce value-added chemicals such methane, ethylene, ethanol, etc. [27] [28] [29] The electrolysis of carbon dioxide gives formate or carbon monoxide, but sometimes more elaborate organic compounds such as ethylene. [30] This technology is under research as a carbon-neutral route to organic compounds. [31] [32]

Electrolysis of acidified water

Electrolysis of water produces hydrogen and oxygen in a ratio of 2 to 1 respectively.

2 H2O(l) → 2 H2(g) + O2(g) E° = +1.229 V

The energy efficiency of water electrolysis varies widely. The efficiency of an electrolyzer is a measure of the enthalpy contained in the hydrogen (to undergo combustion with oxygen or some other later reaction), compared with the input electrical energy. Heat/enthalpy values for hydrogen are well published in science and engineering texts, as 144 MJ/kg. Note that fuel cells (not electrolyzers) cannot use this full amount of heat/enthalpy, which has led to some confusion when calculating efficiency values for both types of technology. In the reaction, some energy is lost as heat. Some reports quote efficiencies between 50% and 70% for alkaline electrolyzers; however, much higher practical efficiencies are available with the use of polymer electrolyte membrane electrolysis and catalytic technology, such as 95% efficiency. [33] [34]

The National Renewable Energy Laboratory estimated in 2006 that 1 kg of hydrogen (roughly equivalent to 3 kg, or 4 liters, of petroleum in energy terms) could be produced by wind powered electrolysis for between US$5.55 in the near term and US$2.27 in the longer term. [35]

About 4% of hydrogen gas produced worldwide is generated by electrolysis, and normally used onsite. Hydrogen is used for the creation of ammonia for fertilizer via the Haber process, and converting heavy petroleum sources to lighter fractions via hydrocracking. Recently, onsite electrolysis has been utilized to capture hydrogen for hydrogen fuel-cells in hydrogen vehicles.

Carbon/hydrocarbon assisted water electrolysis

Recently, to reduce the energy input, the utilization of carbon (coal), alcohols (hydrocarbon solution), and organic solution (glycerol, formic acid, ethylene glycol, etc.) with co-electrolysis of water has been proposed as a viable option. [36] [37] The carbon/hydrocarbon assisted water electrolysis (so-called CAWE) process for hydrogen generation would perform this operation in a single electrochemical reactor. This system energy balance can be required only around 40% electric input with 60% coming from the chemical energy of carbon or hydrocarbon. [38] This process utilizes solid coal/carbon particles or powder as fuels dispersed in acid/alkaline electrolyte in the form of slurry and the carbon contained source co-assist in the electrolysis process as following theoretical overall reactions: [39]

Carbon/Coal slurry (C + 2H2O) → CO2 + 2H2 E′ = 0.21 V (reversible voltage) / E′ = 0.46 V (thermo-neutral voltage)


Carbon/Coal slurry (C + H2O) → CO + H2 E′ = 0.52 V (reversible voltage) / E′ = 0.91 V (thermo-neutral voltage)

Thus, this CAWE approach is that the actual cell overpotential can be significantly reduced to below 1.0 V as compared to 1.5 V for conventional water electrolysis.


A specialized application of electrolysis involves the growth of conductive crystals on one of the electrodes from oxidized or reduced species that are generated in situ. The technique has been used to obtain single crystals of low-dimensional electrical conductors, such as charge-transfer salts and linear chain compounds [40] [41]

See also

Related Research Articles

Electrochemistry Branch of chemistry

Electrochemistry is the branch of physical chemistry concerned with the relationship between electrical potential difference, as a measurable and quantitative phenomenon, and identifiable chemical change, with the potential difference as an outcome of a particular chemical change, or vice versa. These reactions involve electrons moving via an electronically-conducting phase between electrodes separated by an ionically conducting and electronically insulating electrolyte.

Electrochemical cell Electro-chemical device

An electrochemical cell is a device capable of either generating electrical energy from chemical reactions or using electrical energy to cause chemical reactions. The electrochemical cells which generate an electric current are called voltaic or galvanic cells and those that generate chemical reactions, via electrolysis for example, are called electrolytic cells. A common example of a galvanic cell is a standard 1.5 volt cell meant for consumer use. A battery consists of one or more cells, connected in parallel, series or series-and-parallel pattern.

An electrolyte is a medium containing ions that is electrically conducting through the movement of those ions, but not conducting electrons. This includes most soluble salts, acids, and bases dissolved in a polar solvent, such as water. Upon dissolving, the substance separates into cations and anions, which disperse uniformly throughout the solvent. Solid-state electrolytes also exist. In medicine and sometimes in chemistry, the term electrolyte refers to the substance that is dissolved.

The Hall–Héroult process is the major industrial process for smelting aluminium. It involves dissolving aluminium oxide (alumina) in molten cryolite, and electrolyzing the molten salt bath, typically in a purpose-built cell. The Hall–Héroult process applied at industrial scale happens at 940–980 °C and produces 99.5–99.8% pure aluminium. Recycled aluminum requires no electrolysis, thus it does not end up in this process. This process contributes to climate change through the emission of carbon dioxide and fluorocarbons in the electrolytic reaction and consumption of large amounts of electrical energy.

In chemistry, a reducing agent is a chemical species that "donates" an electron to an electron recipient. Examples of substances that are commonly reducing agents include the Earth metals, formic acid, oxalic acid, and sulfite compounds.

Galvanic cell Electrochemical device

A galvanic cell or voltaic cell, named after the scientists Luigi Galvani and Alessandro Volta, respectively, is an electrochemical cell in which an electric current is generated from spontaneous Oxidation-Reduction reactions. A common apparatus generally consists of two different metals, each immersed in separate beakers containing their respective metal ions in solution that are connected by a salt bridge or separated by a porous membrane.

Electrolytic cell Cell that uses electrical energy to drive a non-spontaneous redox reaction

An electrolytic cell is an electrochemical cell that utilizes an external source of electrical energy to drive a chemical reaction that would not otherwise occur. This is in contrast to a galvanic cell, which itself is a source of electrical energy and the foundation of a battery. The net reaction taking place in a galvanic cell is a spontaneous reaction, i.e, the Gibbs free energy remains negative, while the net reaction taking place in an electrolytic cell is the reverse of this spontaneous reaction, i.e, the Gibbs free energy is positive.

The chloralkali process is an industrial process for the electrolysis of sodium chloride solutions. It is the technology used to produce chlorine and sodium hydroxide, which are commodity chemicals required by industry. 35 million tons of chlorine were prepared by this process in 1987. The chlorine and sodium hydroxide produced in this process are widely used in the chemical industry.

Alkaline fuel cell Type of fuel cell

The alkaline fuel cell (AFC), also known as the Bacon fuel cell after its British inventor, Francis Thomas Bacon, is one of the most developed fuel cell technologies. Alkaline fuel cells consume hydrogen and pure oxygen, to produce potable water, heat, and electricity. They are among the most efficient fuel cells, having the potential to reach 70%.

Sodium chlorate Chemical compound

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Electrocoagulation (EC) is a technique used for wastewater treatment, wash water treatment, industrially processed water, and medical treatment. Electrocoagulation has become a rapidly growing area of wastewater treatment due to its ability to remove contaminants that are generally more difficult to remove by filtration or chemical treatment systems, such as emulsified oil, total petroleum hydrocarbons, refractory organics, suspended solids, and heavy metals. There are many brands of electrocoagulation devices available and they can range in complexity from a simple anode and cathode to much more complex devices with control over electrode potentials, passivation, anode consumption, cell REDOX potentials as well as the introduction of ultrasonic sound, ultraviolet light and a range of gases and reactants to achieve so-called Advanced Oxidation Processes for refractory or recalcitrant organic substances.

Electrolysis of water Electricity-induced chemical reaction

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In electrochemistry, overpotential is the potential difference (voltage) between a half-reaction's thermodynamically determined reduction potential and the potential at which the redox event is experimentally observed. The term is directly related to a cell's voltage efficiency. In an electrolytic cell the existence of overpotential implies that the cell requires more energy than thermodynamically expected to drive a reaction. In a galvanic cell the existence of overpotential means less energy is recovered than thermodynamics predicts. In each case the extra/missing energy is lost as heat. The quantity of overpotential is specific to each cell design and varies across cells and operational conditions, even for the same reaction. Overpotential is experimentally determined by measuring the potential at which a given current density is achieved.

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Solid oxide electrolyzer cell Type of fuel cell

A solid oxide electrolyzer cell (SOEC) is a solid oxide fuel cell that runs in regenerative mode to achieve the electrolysis of water by using a solid oxide, or ceramic, electrolyte to produce hydrogen gas and oxygen. The production of pure hydrogen is compelling because it is a clean fuel that can be stored, making it a potential alternative to batteries, methane, and other energy sources. Electrolysis is currently the most promising method of hydrogen production from water due to high efficiency of conversion and relatively low required energy input when compared to thermochemical and photocatalytic methods.

The lithium–air battery (Li–air) is a metal–air electrochemical cell or battery chemistry that uses oxidation of lithium at the anode and reduction of oxygen at the cathode to induce a current flow.

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Electro-oxidation(EO), also known as anodic oxidation or electrochemical oxidation, is a technique used for wastewater treatment, mainly for industrial effluents, and is a type of advanced oxidation process (AOP). The most general layout comprises two electrodes, operating as anode and cathode, connected to a power source. When an energy input and sufficient supporting electrolyte are provided to the system, strong oxidizing species are formed, which interact with the contaminants and degrade them. The refractory compounds are thus converted into reaction intermediates and, ultimately, into water and CO2 by complete mineralization.


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