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Illustration of a Hofmann electrolysis apparatus used in a school laboratory
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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."

Etymology

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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.

History

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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 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 that 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]

In 1902 Polish engineer and inventor Stanisław Łaszczyński filed for and obtained Polish patent for the electrolysis of copper and zinc.[16][17][18]

Timeline

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Overview

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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:

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.[24] 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

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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 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:[25]

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.[26]

Decomposition potential

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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.[27]

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

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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:

Fe2+
(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:

Fe(CN)3-
6 + e− → Fe(CN)4-
6

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

+ 2 e− + 2 H+ →

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

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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

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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.[28] Electrolysis of ethanol with pulsed current evolves an aldehyde instead of primarily an acid.[29]

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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

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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).[30][31]

Basic membrane cell used in the electrolysis of brine. At the anode (A), chloride (Cl−) is oxidized to chlorine. The ion-selective membrane (B) allows the counterion Na+ to freely flow across, but prevents anions such as hydroxide (OH−) and chloride from diffusing across. At the cathode (C), water is reduced to hydroxide and hydrogen gas. The net process is the electrolysis of an aqueous solution of NaCl into industrially useful products sodium hydroxide (NaOH) and chlorine gas.

Electrofluorination

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Many organofluorine compounds are produced by electrofluorination.[32] 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

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Acrylonitrile is converted to adiponitrile on an industrial scale via electrocatalysis.[33]

Hall–Héroult process for producing aluminium

Electroplating and electrowinning processes

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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)

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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

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Competing half-reactions in solution electrolysis

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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.

Half-reaction E° (V) Ref.
Na+ + e− ⇌ Na(s) −2.71 [34]
Zn2+ + 2 e− ⇌ Zn(s) −0.7618 [35]
2 H+ + 2 e− ⇌ H2(g) ≡ 0 [35]
Br2(aq) + 2 e− ⇌ 2 Br− +1.0873 [35]
O2(g) + 4 H+ + 4 e− ⇌ 2 H2O +1.23 [34]
Cl2(g) + 2 e− ⇌ 2 Cl− +1.36 [34]
S
2O2−
8 + 2 e− ⇌ 2 SO2−
4 		+2.07 [34]

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):

  • the electrode potential for the reduction producing hydrogen is −0.41 V,
  • the electrode potential for the oxidation producing oxygen is +0.82 V.

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.[36] 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.

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Electrolysis of carbon dioxide

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Main article: Electrochemical reduction of carbon dioxide

The electrochemical reduction or electrocatalytic conversion of CO2 can produce value-added chemicals such as methane, ethylene, ethanol, etc.[37][38][39] The electrolysis of carbon dioxide gives formate or carbon monoxide, but sometimes more elaborate organic compounds such as ethylene.[40] This technology is under research as a carbon-neutral route to organic compounds.[41][42]

Electrolysis of acidified water

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Main article: Electrolysis of 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 electrolyser 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 (40 kWh/kg). Note that fuel cells (not electrolysers) 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 electrolysers (50 kWh/kg);[43] however, higher practical efficiencies are available with the use of polymer electrolyte membrane electrolysis and catalytic technology, such as 95% efficiency.[44][45]

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.[46]

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. Onsite electrolysis has been utilized to capture hydrogen for hydrogen fuel-cells in hydrogen vehicles.

Carbon/hydrocarbon assisted water electrolysis

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Main article: Hydrogen production

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.[47][48] 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.[49] 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:[50]

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

or

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.

Electrocrystallization

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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.[51][52]

Electrolysis of Iron Ore

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The current method of producing steel from iron ore is very carbon intensive, in part to the direct release of CO2 in the blast furnace. A study of steel making in Germany found that producing 1 ton of steel emitted 2.1 tons of CO2e with 22% of that being direct emissions from the blast furnace.[53] As of 2022, steel production contributes 7–9% of global emissions.[54] Electrolysis of iron can eliminate direct emissions and further reduce emissions if the electricity is created from green energy.

The small-scale electrolysis of iron has been successfully reported by dissolving it in molten oxide salts and using a platinum anode.[55] Oxygen anions form oxygen gas and electrons at the anode. Iron cations consume electrons and form iron metal at the cathode. This method was performed a temperature of 1550 °C which presents a significant challenge to maintaining the reaction. Particularly, anode corrosion is a concern at these temperatures.

Additionally, the low temperature reduction of iron oxide by dissolving it in alkaline water has been reported.[56] The temperature is much lower than traditional iron production at 114 °C. The low temperatures also tend to correlate with higher current efficiencies, with an efficiency of 95% being reported. While these methods are promising, they struggle to be cost competitive because of the large economies of scale keeping the price of blast furnace iron low.

Electrolysis of seawater

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A 2020 study investigated direct electrolysis of seawater, alkaline electrolysis, proton-exchange membrane electrolysis, and solid oxide electrolysis.[57] Direct electrolysis of seawater follows known processes, forming an electrolysis cell in which the seawater acts as the electrolyte to allow for the reaction at the anode, 2 Cl−(aq) → Cl2(g) + 2e− and the reaction at the cathode, 2 H2O(l) + 2 e− → H2(g) + 2OH−(aq). The inclusion of magnesium and calcium ions in the seawater makes the production of alkali hydroxides possible that could form scales in the electrolyser cell, cutting down on lifespan and increasing the need for maintenance. The alkaline electrolysers operate with the following reactions at the anode, 2 OH−(aq) → 1/2 O2(g) + H2O(l) + 2 e−, and at the cathode, 2 H2O(l) + 2 e− → H2(g) + 2 OH−(aq), and use high base solutions as electrolytes, operating at 60–90 °C (140–194 °F) and need additional separators to ensure the gas phase hydrogen and oxygen remain separate. The electrolyte can easily get contaminated, but the alkaline electrolyser can operate under pressure to improve energy consumption. The electrodes can be made of inexpensive materials and there's no requirement for an expensive catalyst in the design.

Many alternatives to this simple electrolyzer described exist. Micro-electrolyzer designs are able to eliminate the separator requirement by designing the internal flow to separate the gases autonomously. See for example US 12116679B2  ("Device and method for large scale harvesting of solar energy through hydrogen production") where the operating pressure is increased to the point of Chlorine liquefaction so that sea water electrolyzer can proceed in a locally alkaline electrolytic fluid. Removing separators allows operating at very high temperatures. The structural design allows for operations at upto 700 bar thereby eliminating the need for Hydrogen compressors.

Proton-exchange membrane electrolysers operate with the reactions at the anode, H2O(l) → 1/2 O2(g) + 2 H+(aq) + 2 e− and cathode, 2 H+(aq) + 2 e− → H2(g), at temperatures of 60–80 °C (140–176 °F), using a solid polymer electrolyte and requiring higher costs of processing to allow the solid electrolyte to touch uniformly to the electrodes. Similar to the alkaline electrolyser, the proton exchange membrane electrolyser can operate at higher pressures, reducing the energy costs required to compress the hydrogen gas afterward, but the proton exchange membrane electrolyser also benefits from rapid response times to changes in power requirements or demands and not needing maintenance, at the cost of having a faster inherent degradation rate and being the most vulnerable to impurities in the water.

Solid oxide electrolysers run the reactions O−2(g) → 1/2 O2(g) + 2 e− at the anode and H2O(g) + 2 e− → H2(g) + O−2(g) at the cathode. The solid oxide electrolysers require high temperatures (700–1,000 °C (1,292–1,832 °F)) to operate, generating superheated steam. They suffer from degradation when turned off, making it a more inflexible hydrogen generation technology. In a selected series of multiple-criteria decision-analysis comparisons in which the highest priority was placed on economic operation costs followed equally by environmental and social criteria, it was found that the proton exchange membrane electrolyser offered the most suitable combination of values (e.g., investment cost, maintenance, and operation cost, resistance to impurities, specific energy for hydrogen production at sea, risk of environmental impact, etc.), followed by the alkaline electrolyser, with the alkaline electrolyser being the most economically feasible, but more hazardous in terms of safety and environmental concerns due to the need for basic electrolyte solutions as opposed to the solid polymers used in proton-exchange membranes. Due to the methods conducted in multiple-criteria decision analysis, non-objective weights are applied to the various factors, and so multiple methods of decision analysis were performed simultaneously to examine the electrolysers in a way that minimizes the effects of bias on the performance conclusions.

See also

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References

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Revisions and contributorsEdit on WikipediaRead on Wikipedia

Electrolysis

View on Grokipedia
from Grokipedia
Electrolysis is a chemical process in which an drives a non-spontaneous reaction, typically decomposing an into its constituent parts such as elements or simpler compounds. This occurs in an where an external voltage overcomes the positive change of the reaction, with oxidation at the and reduction at the . The discovery of electrolysis followed Alessandro Volta's invention of the in 1800, enabling William Nicholson and Anthony Carlisle to decompose into and oxygen gases. In 1833, formulated the quantitative laws of electrolysis, establishing that the mass of a substance altered at an is directly proportional to the of transferred and that the amounts of different substances deposited or liberated by a fixed of are proportional to their equivalent weights. These laws provided a foundational framework for , linking electrical charge to chemical change through the , approximately 96,485 coulombs per mole of electrons. Faraday's work built on earlier demonstrations by , who isolated several and alkaline earth metals via electrolysis of molten salts. Electrolysis finds extensive industrial applications, including the chlor-alkali process for producing and , the Hall-Héroult process for aluminum extraction from , and electrorefining of metals like for high purity. More recently, it has gained prominence in water electrolysis for production, essential for decarbonizing sectors such as and transportation, though challenges persist in achieving cost-effective efficiency at scale.

Etymology and Historical Context

Etymology

The term "electrolysis" derives from the Greek words ἤλεκτρον (ḗlektron), meaning "amber"—the substance from which static electricity was first generated via friction in antiquity—and λύσις (lúsis), denoting "a loosening," "dissolution," or "breaking apart." Michael Faraday introduced the word in 1834 to precisely capture the action of an electric current in separating or decomposing electrolytic substances into their constituent elements, emphasizing its descriptive fidelity over prior magnetic-inspired terminology like "poles." Before this coinage, the phenomenon was commonly referred to as "electrical decomposition" or "electrochemical decomposition" in scientific literature.

Early Experiments and Discoveries

In 1789, Dutch physicists Adriaan Paets van Troostwijk and Jan Rudolph Deiman performed experiments using high-voltage electrostatic generators to pass electric sparks through water, observing the decomposition into two distinct gases later identified as and oxygen. Their work, published in the proceedings of the Hollandsche Maatschappij der Wetenschappen, represented an early empirical demonstration of via electrical means, though limited by the intermittent nature of . The development of Alessandro Volta's in late 1799, publicly announced in 1800, introduced a device capable of producing a steady from stacked electrochemical cells of and disks separated by brine-soaked cardboard. This innovation enabled sustained electrolysis experiments; within months, British scientists William Nicholson and Anthony Carlisle used the pile to decompose into and oxygen bubbles at electrodes immersed in the liquid, confirming the process's reproducibility with galvanic . Building on these advances, conducted electrolysis of molten salts in the early 1800s at the Royal Institution, isolating elemental sodium on October 6, 1807, by applying current from large voltaic batteries to heated , and potassium shortly thereafter from molten on the same day. These experiments yielded shiny, reactive metal globules that ignited in air, providing the first isolation of alkali metals and highlighting electrolysis's potential for extracting elements from compounds previously resistant to chemical reduction.

Key Theoretical and Practical Milestones

Michael Faraday established the foundational quantitative principles of electrolysis through his two laws, formulated based on experiments conducted between 1831 and 1833 and published in 1833. The first law states that the mass of a substance deposited or liberated at an electrode is directly proportional to the total electric charge passed through the electrolyte, expressed as m=QF⋅Mzm = \frac{Q}{F} \cdot \frac{M}{z}, where mm is mass, QQ is charge, FF is the Faraday constant, MM is molar mass, and zz is the number of electrons transferred. The second law asserts that for a fixed quantity of charge, the masses of different substances produced are proportional to their equivalent weights, linking electrolysis to stoichiometry and confirming that one Faraday of charge (approximately 96,485 coulombs) liberates one equivalent of substance, such as 1 gram-equivalent of hydrogen or oxygen gas. These laws provided the empirical basis for predicting electrolytic yields, enabling precise control over gas production volumes, as demonstrated in subsequent voltameter setups where hydrogen and oxygen volumes adhered to a 2:1 ratio by Faraday's equivalence. In 1905, Julius Tafel introduced the concept of through his empirical equation, η=a+blog⁡i\eta = a + b \log i, which quantifies the additional voltage required beyond the thermodynamic minimum to drive reactions at significant rates due to kinetic barriers at the surface. This relation, derived from evolution studies on various metals, revealed that overpotentials arise primarily from hurdles, with the Tafel slope bb (often 0.12 V/decade for ) indicating sensitivity to material. Practically, Tafel's work explained why cell voltages in electrolysis exceed theoretical values by 0.3-1 V or more, guiding selection and development to minimize losses, thus marking a shift from purely thermodynamic to kinetic understanding essential for scalable processes. The 1920s and 1930s saw the commercialization of chloralkali electrolysis, with diaphragm cells dominating new installations for separating anode and cathode compartments using porous barriers to prevent formation, enabling efficient production of gas and from . Mercury cells, invented earlier but scaled commercially during this period, employed a flowing mercury to form a , reducing evolution and yielding purer caustic soda, as exemplified by operational plants from the mid-1930s onward. These configurations achieved industrial capacities exceeding thousands of tons annually, driven by demand for chemicals in wartime and postwar economies, though mercury cells introduced environmental concerns later recognized. Post-World War II advancements included the development of solid electrolytes for electrolysis, with General Electric's introduction of solid polymer electrolyte (SPE) technology in the early using cation-exchange membranes like Nafion precursors to eliminate liquid electrolytes, improving safety and efficiency in . High-temperature variants emerged concurrently, leveraging ceramic solid electrolytes such as operating above 800°C to reduce electrical energy input by integrating process heat, as initial concepts for steam electrolysis demonstrated lower overpotentials through enhanced ion mobility. These innovations laid groundwork for compact, durable electrolyzers, though challenges in material stability persisted until refined in subsequent decades.

Industrial and Scientific Evolution

The commercialization of the Hall-Héroult process in the late 1880s marked the onset of large-scale industrial electrolysis, transforming aluminum from a rare metal into a viable industrial material through electrolytic reduction of alumina dissolved in molten cryolite. demonstrated the process in 1886, leading to the formation of the Pittsburgh Reduction Company in 1888, which achieved initial commercial production and scaled to several tons annually by 1893, driven by demand for lightweight materials in emerging electrical and transportation sectors. World War I demands for hydrogen to inflate observation balloons and dirigibles spurred advancements in water electrolysis, as military needs for on-site, scalable gas production necessitated robust, high-volume electrolyzers operating continuously under variable loads. This wartime pressure refined electrode materials and cell designs, laying groundwork for post-war efficiency gains amid energy constraints. In the chlor-alkali industry, the introduction of cells in the late represented a pivotal adaptation, replacing mercury and diaphragm cells with selective barriers that minimized product mixing and reduced by more than 30% relative to predecessors, primarily through lower cell voltages and improved current efficiencies. The 1970s oil crises amplified economic drivers for process optimization, accelerating the shift toward continuous operations in electrolyzers, which supplanted batch-like early configurations by enabling steady-state production and current densities rising from approximately 0.1 A/cm² in initial industrial setups to over 1 A/cm² in optimized systems, thereby boosting productivity while curbing operational costs.

Fundamental Principles

Core Process of Electrolysis

Electrolysis occurs when a direct current voltage is applied across two electrodes immersed in an electrolyte, such as an aqueous solution or molten salt, causing cations to migrate toward the negatively charged cathode and anions toward the positively charged anode. This ion migration completes the internal circuit, facilitating the flow of electrons externally from anode to cathode. At the cathode, reduction reactions take place, where electrons are gained by species in the electrolyte, often leading to gas evolution such as hydrogen in water-based systems. At the anode, oxidation reactions occur, involving electron loss and typically producing gases like oxygen. The electrolyte's role is to provide dissociated ions that enhance solution conductivity, enabling efficient charge transfer without which pure water would exhibit insufficient ionic mobility for practical current passage. In the electrolysis of water, the overall decomposition reaction is 2H2O→2H2+O22H_2O \rightarrow 2H_2 + O_2, driven by the applied voltage. The cathodic half-reaction in neutral or basic conditions is 2H2O+2e−→H2+2OH−2H_2O + 2e^- \rightarrow H_2 + 2OH^-, while the anodic half-reaction is 4OH−→O2+2H2O+4e−4OH^- \rightarrow O_2 + 2H_2O + 4e^-; in acidic media, these adjust to involve H+H^+ ions accordingly. Electrolytes like sulfuric acid or potassium hydroxide are commonly used to increase ion concentration and thus conductivity in aqueous water electrolysis setups.

Thermodynamic Foundations

Electrolysis constitutes a non-spontaneous electrochemical process driven by an externally applied voltage that supplies the minimum work equivalent to the change (ΔG) of the reaction. For the splitting of , 2H₂O(l) → 2H₂(g) + O₂(g), the standard ΔG° equals 237 kJ/mol at 25°C and 1 , reflecting the energy barrier against spontaneity under these conditions. This ΔG° determines the reversible cell voltage via the relation ΔG = -nFE_rev, where n = 2 mol of electrons transferred per mol of decomposed and F = 96,485 C/mol is the , yielding E_rev = 1.23 V as the theoretical minimum voltage for the process. The reversible voltage differs from the total change ΔH° = 285.8 kJ/mol for the reaction, as ΔG° = ΔH° - TΔS°, with the term TΔS° ≈ 48.6 kJ/mol at 298 K arising from the increased disorder of gaseous products relative to liquid . In a reversible isothermal electrolysis, electrical input covers ΔG°, while TΔS° is absorbed from the bath, enabling the higher heating value of produced (corresponding to ΔH°) to exceed the consumed, thus permitting thermodynamic efficiencies above 100% when referenced to electrical input alone. The thermoneutral voltage, ΔH°/nF ≈ 1.48 V, marks the point of zero net exchange; operation between 1.23 V and 1.48 V requires input, while exceeding 1.48 V generates excess . This reversible potential varies with operating conditions per thermodynamic principles. Temperature elevation reduces E_rev, as the -TΔS° term diminishes ΔG° for the endothermic, entropy-increasing reaction. increases E_rev via the Nernst equation's logarithmic dependence on partial pressures of gaseous products, favoring higher voltages under compression. concentration affects activities, further modulating the potential through Nernstian shifts, though standard conditions assume pure water or dilute solutions. These dependencies underscore electrolysis's inherent thermodynamic constraints, independent of kinetic barriers.

Electrode Reactions and Overpotentials

In water electrolysis, the cathodic electrode reaction primarily involves the (HER), where in acidic conditions, protons are reduced according to 2H⁺ + 2e⁻ → H₂ with a standard reduction potential of 0 V versus the (SHE) at 25°C. In alkaline media, the reaction shifts to 2H₂O + 2e⁻ → H₂ + 2OH⁻, with an equilibrium potential of approximately -0.828 V versus SHE at 14. At the anode, the reaction (OER) occurs, represented in acidic solution as 2H₂O → O₂ + 4H⁺ + 4e⁻ with a standard potential of 1.229 V versus SHE, yielding a theoretical cell voltage of 1.229 V under standard conditions (298 K, 1 atm, 0). In basic conditions, the anodic reaction is 4OH⁻ → O₂ + 2H₂O + 4e⁻, adjusted to about 0.401 V versus SHE at 14. These half-cell reactions dictate the ideal reversible potential, but empirical cell voltages exceed this due to overpotentials that impose additional voltage requirements. Overpotentials arise from kinetic, , and resistive barriers, deviating from thermodynamic ideals. Activation overpotential stems from the energy barrier for charge transfer at the electrode-electrolyte interface, particularly pronounced for multi-step OER involving oxygen intermediates, unlike the relatively facile two-electron HER. This is quantified by the Butler-Volmer equation, which relates net ii to η\eta as i=i0[exp⁡(αanFηRT)−exp⁡(−αcnFηRT)]i = i_0 \left[ \exp\left(\frac{\alpha_a n F \eta}{RT}\right) - \exp\left(-\frac{\alpha_c n F \eta}{RT}\right) \right], where i0i_0 is the , α\alpha are transfer coefficients, nn is the number of electrons, FF is Faraday's constant, RR is the , and TT is ; at high overpotentials, it approximates the η=a+blog⁡∣i∣\eta = a + b \log|i|, with Tafel slope b=2.303RTαnFb = \frac{2.303 RT}{\alpha n F}. Concentration overpotential results from limitations, causing reactant depletion or product accumulation near the electrode, especially at high current densities, while ohmic overpotential reflects voltage drops from ionic resistance in the and electronic resistance in components, following ηohmic=iR\eta_{ohmic} = i R, where RR is the total resistance. Electrode material selection critically influences magnitudes, with exhibiting low activation overpotentials for HER due to its high exchange current density and optimal binding energy for hydrogen intermediates, often requiring less than 50 mV overpotential at 10 mA/cm² in acidic media. However, 's efficacy diminishes in alkaline environments, and for OER, it performs poorly compared to oxides like or , which lower barriers but suffer in acidic or oxidative conditions, leading to degradation over time. Empirical studies confirm that non-precious alternatives, such as nickel-based catalysts, reduce costs but incur higher activation losses, underscoring trade-offs between activity and stability in practical electrolysis. These dependencies highlight how surface properties and composition causally amplify deviations from standard potentials.

Quantitative Metrics and Efficiency Definitions

Faradaic efficiency, also known as current efficiency, quantifies the selectivity of an electrolysis process by measuring the fraction of electrical charge that contributes to the desired electrode reaction rather than side reactions. It is defined as the ratio of the actual amount of product formed (in moles) to the theoretical amount predicted by Faraday's laws based on the total charge passed, expressed as ηF=nactualntheoretical=nactual⋅z⋅FQ\eta_F = \frac{n_{\text{actual}}}{n_{\text{theoretical}}} = \frac{n_{\text{actual}} \cdot z \cdot F}{Q}, where nn is moles of product, zz is electrons transferred per , FF is Faraday's constant, and QQ is total charge. Values approaching 100% indicate minimal parasitic currents, though real systems often exhibit losses from competing reactions like gas evolution or recombination. Voltage efficiency captures the thermodynamic losses in the cell potential, defined as ηV=EthermoEcell\eta_V = \frac{E_{\text{thermo}}}{E_{\text{cell}}}, where EthermoE_{\text{thermo}} is the theoretical reversible voltage (e.g., 1.23 V for water electrolysis at standard conditions) and EcellE_{\text{cell}} is the actual operating voltage including overpotentials. Some definitions use the thermoneutral voltage (approximately 1.48 V for water electrolysis, accounting for reaction enthalpy) to reflect heat generation, emphasizing that efficiencies above 100% electrical input are possible with external heat but require careful distinction from pure electrical metrics. This metric highlights overpotentials from kinetics, mass transport, and ohmic resistance, which degrade performance independently of Faradaic losses. Overall energy efficiency for electrolysis, particularly hydrogen production, integrates these factors as η=ηF⋅ηV⋅ηother\eta = \eta_F \cdot \eta_V \cdot \eta_{\text{other}}, but is commonly computed as the ratio of the hydrogen's heating value to total electrical energy input: η=HHV (or LHV)⋅mH2Einput\eta = \frac{\text{HHV (or LHV)} \cdot m_{\text{H}_2}}{E_{\text{input}}}. The higher heating value (HHV, 39.4 kWh/kg H₂) includes the latent heat of water vapor condensation, yielding conservative efficiencies, while the lower heating value (LHV, 33.3 kWh/kg H₂) assumes gaseous exhaust and reports higher percentages for the same input. Low-temperature electrolyzers (e.g., alkaline or PEM) typically achieve 60-80% efficiency on an HHV basis at commercial scales, limited by electrical-only inputs and waste heat dissipation. Total thermal efficiency, incorporating recoverable waste heat, can exceed 90% in high-temperature solid oxide electrolysis cells (SOEC) when coupled with industrial heat sources, as the endothermic reaction leverages external thermal energy to reduce electrical demand. This distinction debunks claims of inherent low efficiency by revealing that simplistic electrical metrics overlook viable heat integration for causal energy balances.

Practical Configurations and Variations

Types of Electrolyzers

Alkaline electrolyzers, the most established type, employ a liquid electrolyte typically consisting of 20-40% (KOH) in , separated by a porous diaphragm such as or polymeric separators to minimize gas crossover. They operate at ambient to moderate temperatures (60-90°C) and , achieving electrical efficiencies of 60-70% based on the higher heating value (HHV) of . Current densities are generally limited to 0.2-0.5 A/cm², supporting large-scale with stack lifetimes exceeding 80,000 hours, though startup times range from minutes to hours due to the liquid electrolyte's thermal inertia. Hydrogen purity is around 99.5-99.9%, often requiring downstream purification to remove trace oxygen. These systems trade higher capital costs for mature reliability but face challenges in dynamic load response compared to solid-electrolyte variants. Proton exchange membrane (PEM) electrolyzers utilize a solid polymer electrolyte membrane, such as sulfonated tetrafluoroethylene-based , which conducts protons while preventing gas mixing. Operating at 50-80°C and capable of current densities up to 2 A/cm², they deliver purities exceeding 99.99% without additional separation, with efficiencies of 65-80% HHV. Startup occurs in seconds, enabling rapid response to intermittent power inputs, and compact designs facilitate modular deployment. However, reliance on precious metal catalysts ( for evolution, for oxygen evolution) elevates costs, with stack lifetimes around 40,000-80,000 hours under continuous operation. Scalability is constrained by durability and catalyst loading, positioning PEM for applications prioritizing purity and responsiveness over lowest upfront expense. Anion exchange membrane (AEM) electrolyzers represent an emerging hybrid, featuring a solid that enables transport in alkaline conditions. They operate at 40-80°C with current densities approaching 1 A/cm² in prototypes, potentially achieving efficiencies similar to alkaline systems (60-70% HHV) while offering PEM-like compactness and startup times under a minute. purity matches PEM levels (>99.9%), and the technology avoids precious metals by using non-noble catalysts compatible with alkaline media. Development since the has progressed to pilot scales, but stability and conductivity remain barriers, limiting commercial scalability and stack lifetimes to below 10,000 hours currently. AEM variants promise cost reductions through material substitutions but require further validation for high-pressure operations. Solid oxide electrolyzers (SOEC) employ a ceramic electrolyte, often (YSZ), functioning at elevated temperatures of 700-900°C to enhance kinetics and enable co-electrolysis with or CO₂. Efficiencies reach 80-90% HHV when integrating , surpassing low-temperature types by leveraging to reduce electrical input, with current densities up to 1-2 A/cm² feasible. Hydrogen purity exceeds 99.9%, and the high-temperature operation supports process integration but demands slow startups (hours) and robust materials to withstand thermal cycling, resulting in stack lifetimes of 5,000-20,000 hours. Scalability is emerging via pressurized designs, though material degradation from oxidation and poses trade-offs against efficiency gains. electrolyzers, variants using ionic liquids or fused salts at 200-600°C, offer similar high-efficiency potential for specialized non-aqueous processes but see limited adoption for pure due to and complexity.

Operational Parameters and Conditions

, defined as the per unit area (typically in A/cm²), directly influences the and production yield in electrolysis systems by driving the kinetics of at the electrodes. Higher current densities accelerate gas evolution and output rates, enabling smaller electrode areas and reduced capital costs, but they also amplify ohmic losses and mass transport limitations, potentially decreasing faradaic efficiency if not managed. In advanced alkaline electrolyzers, nominal current densities of up to 1.8 A/cm² have been demonstrated as achievable under optimized conditions. Temperature affects ionic conductivity, reaction kinetics, and gas solubility, with elevated values generally enhancing charge transfer rates and reducing activation overpotentials to improve yields. For , operating temperatures in the range of 60–100 °C promote faster kinetics, though exceeding 80 °C risks accelerated of components like electrodes and diaphragms. An upper limit of 80 °C is often set to balance efficiency gains against material degradation. Operating pressure modulates gas bubble behavior and solubility, influencing and crossover rates, with higher pressures up to 30 bar facilitating downstream gas handling by minimizing subsequent mechanical compression energy needs. However, pressures above 8 bar can constrain minimum load flexibility by exacerbating diffusive limitations at low currents. Elevated pressures also tend to increase cell voltage requirements due to altered , though they can suppress certain parasitic reactions. Electrolyte concentration governs solution conductivity and mobility, thereby impacting overall cell resistance and current distribution uniformity, which in turn affect yields by minimizing uneven reaction kinetics. In alkaline systems, concentrations of 20–40 wt% KOH optimize conductivity for efficient operation, as higher levels enhance ion availability and reduce ohmic drops, but excessive concentrations promote and viscosity increases that hinder . Lower concentrations suffice for conductivity in some setups but may elevate resistance losses. Electrolyte flow rates influence convective mass transport and bubble detachment, critical for sustaining uniform kinetics and high yields by preventing coverage of active sites that would otherwise induce diffusion limitations. Adequate flow rates, such as 1.1 L/min in certain porous configurations, ensure effective bubble removal without significant voltage penalties, while optimized rates under elevated pressures further curb bubble growth to preserve . Insufficient flows lead to aggregation and retention, exacerbating local overpotentials.

Competing Reactions and Selectivity Issues

In aqueous electrolysis, competing anodic reactions such as (OER: 2H2O→O2+[4H](/page/4−H)++4e−2H_2O \rightarrow O_2 + [4H](/page/4-H)^+ + 4e^-) versus oxidation (ClER: 2Cl−→Cl2+2e−2Cl^- \rightarrow Cl_2 + 2e^-) diminish selectivity, particularly in saline electrolytes like seawater. Thermodynamically, the standard reversible potential for OER is 1.23 V versus SHE at pH 0, while ClER is 1.36 V, but Pourbaix diagrams illustrate pH-dependent stability regions where alkaline conditions favor OER by shifting oxidation toward (HOCl/ClO−^-) formation, increasing its to over 1.5 V. Kinetically, OER demands high overpotentials (often 300–500 mV at 10 mA/cm² on catalysts), whereas ClER exhibits lower barriers on many surfaces, resulting in yields exceeding 50% without intervention in neutral-to-acidic media. Cathodic selectivity faces analogous challenges, with hydrogen evolution (HER: 2H++2e−→H22H^+ + 2e^- \rightarrow H_2) competing against metal reduction (e.g., Mn++ne−→MM^{n+} + ne^- \rightarrow M) in solutions containing electroactive cations. For metals like (E0=−0.76E^0 = -0.76 V vs. SHE) or aluminum, whose reduction potentials lie negative to HER's 0 V, proton or reduction dominates if the favors fast HER kinetics, leading to current efficiencies below 70% in some processes. exemplifies overpotential-driven bias, requiring only ~20 mV for HER at 10 mA/cm² due to optimal H adsorption energetics, but over 400 mV for OER stemming from unfavorable OOH* intermediates, inherently skewing toward over oxygen or deposition products. Strategies to enhance selectivity leverage catalyst design and pH modulation, though each incurs kinetic penalties. Selective OER catalysts, such as polymorphic MnO₂, lower OER overpotentials relative to ClER by stabilizing peroxo intermediates, achieving >95% oxygen faradaic efficiency in chloride media at currents up to 100 mA/cm². Alkaline (>13) thermodynamically widens the OER-ClER gap via Nernst shifts but can elevate HER overpotentials on non-precious metals, necessitating additional voltage (up to 0.2 V extra) and raising energy costs by 10–20%. For cathodic deposition, additives or alloyed electrodes increase HER overpotentials (e.g., by poisoning H adsorption sites), yet this often reduces deposition rates, trading yield for purity at the expense of higher applied potentials. These mitigations underscore inherent causal trade-offs, where selectivity gains typically amplify overpotentials elsewhere, constraining overall process viability.

Industrial Applications

Chemical Production Processes

Electrolysis serves as a key method for industrial chemical production, enabling the synthesis of inorganic and organic compounds through controlled electrochemical reactions at scalable volumes. In the , aqueous () undergoes electrolysis to yield gas (Cl₂), (NaOH), and gas (H₂), with the overall reaction 2NaCl + 2H₂O → 2NaOH + Cl₂ + H₂. This process accounts for approximately 90% of global , exceeding 70 million metric tons annually as of 2020. cells, utilizing ion-exchange membranes to separate and compartments, became predominant in the 1980s, replacing mercury and diaphragm cells due to improved energy efficiency and reduced environmental hazards. These cells achieve current efficiencies above 95% for chlorine evolution, with cell voltages typically around 3.0-3.5 V at current densities of 2-4 kA/m². The Simons electrofluorination process, developed in the 1940s, produces perfluorinated organic compounds by direct anodic fluorination of hydrocarbons or derivatives in anhydrous (HF). In this method, the substrate undergoes stepwise replacement of hydrogen atoms with fluorine, yielding perfluorocarbons used in applications like refrigerants and , with the general reaction RF-H + F⁻ → RF-F + H⁻ (simplified at the ). Industrial implementation by Corporation scaled production of compounds such as (PFOA), though phased out in many regions due to persistence concerns; the process operates at potentials of 5-8 V and temperatures of 0-10°C to manage HF's corrosivity. Despite challenges like low yields (often <50%) from side reactions, it remains a unique electrochemical route for fluorination not feasible via chemical means. Electrohydrodimerization of acrylonitrile to adiponitrile, a precursor for nylon-6,6 via hydrogenation to hexamethylenediamine, exemplifies organic electrosynthesis at industrial scale. The cathodic reaction in aqueous emulsion or divided cells couples two acrylonitrile molecules: 2CH₂=CHCN + 2H⁺ + 2e⁻ → NC(CH₂)₄CN, conducted at lead cathodes with potentials around -1.5 to -2.0 V vs. SHE and current densities up to 100 mA/cm². Monsanto commercialized this process in the 1960s, achieving selectivities over 90% and capacities reaching thousands of tons per year before transition to ; BASF later operated similar facilities with energy consumptions of about 2.5-3.0 kWh/kg adiponitrile. This dimerization highlights electrolysis's advantage in selective C-C bond formation under mild conditions, avoiding high-pressure catalysis.

Metallurgical Applications

Electrolysis plays a central role in metallurgical extraction and refining, particularly for non-ferrous metals where high purity is essential for commercial value. Processes like electrowinning recover metals from aqueous solutions derived from ore leaching, enabling efficient separation and deposition at cathodes. These methods are economically viable for high-value metals such as copper, zinc, and aluminum, as the premium on purity offsets energy costs, with global production relying heavily on electrolytic routes. The Hall-Héroult process, developed in 1886, remains the dominant method for primary aluminum production, electrolyzing alumina (Al₂O₃) dissolved in molten cryolite (Na₃AlF₆) at 950–980°C. Carbon anodes oxidize to CO₂, while aluminum collects at the cathode, requiring cell voltages of 4–5 V to overcome overpotentials and ohmic losses. Commercial energy consumption averages 13–15 kWh per kg of aluminum, far exceeding the theoretical minimum of approximately 6.3 kWh/kg due to inefficiencies like anode effects and heat losses. Electrowinning of copper and zinc from sulfate electrolytes follows similar principles, depositing metals at cathodes from solutions produced by hydrometallurgical leaching of ores or concentrates. For copper, electrowinning from CuSO₄-H₂SO₄ solutions at 1.8–2.5 V yields high-purity cathodes (99.99% Cu) integral to refining, with current densities up to 300 A/m². Zinc electrowinning operates at 3–3.5 V from purified ZnSO₄ solutions, producing slabs for alloying, with energy demands around 3–3.5 kWh/kg Zn. These processes ensure metal recovery exceeding 90% efficiency, critical for economic extraction from low-grade ores. Electroplating applies thin metallic coatings via electrolysis to enhance corrosion resistance and durability of base metals. Nickel plating deposits layers 10–50 μm thick, providing a barrier against oxidation and wear, often on steel or copper substrates. Chromium plating, typically 0.5–1 μm thick over nickel underlayers, offers superior hardness and chemical inertness, with hexavalent chromium baths operating at 50–60°C and 200–400 A/m². These coatings extend service life in harsh environments, justifying their use despite environmental concerns over bath chemistries. Electrochemical machining (ECM) employs anodic dissolution for precision shaping of hard alloys, avoiding mechanical tool wear and heat-affected zones. A shaped cathode tool advances toward the anode workpiece in an electrolyte (e.g., NaCl or NaNO₃ solution) under 10–30 V, removing material at rates up to 0.5 mm/min with tolerances below 0.01 mm. This non-contact process suits complex geometries in superalloys like , used in aerospace components, where traditional machining falters.

Hydrogen Generation and Storage

Water electrolysis serves as a method for generating hydrogen gas (H₂) from water, primarily using electricity to drive the reaction 2H₂O → 2H₂ + O₂, with the produced H₂ potentially serving as an energy storage medium when excess electricity is available. Dedicated electrolyzers for H₂ production currently account for less than 1% of global H₂ output, which totals around 95 million tonnes annually, while steam methane reforming (SMR) of natural gas dominates at approximately 75-95% of production depending on regional data. This limited share reflects electrolysis's higher energy intensity compared to SMR, requiring 50-60 kWh of electricity per kg of H₂ produced versus SMR's ~40-50 kWh thermal equivalent from natural gas, excluding upstream methane emissions. Hydrogen from electrolysis offers high purity levels, typically 99.5-99.9%, with minimal contaminants like oxygen or water vapor, making it suitable for direct use in proton-exchange membrane (PEM) fuel cells without extensive purification, unlike SMR-derived H₂ which requires additional processing to remove CO, CO₂, and sulfur compounds. However, when employed for energy storage—converting electricity to H₂ and later reconverting via fuel cells—the round-trip efficiency remains below 40%, combining electrolysis efficiencies of 70-80% (higher heating value basis) with fuel cell efficiencies of 40-60%, further reduced by compression, storage losses, and system auxiliaries. This thermodynamic penalty arises from the endothermic nature of electrolysis (ΔG ≈ 237 kJ/mol at standard conditions) and entropy losses in reversal, limiting its competitiveness against direct battery storage for short-duration applications. Economic viability hinges on capital expenditures (Capex) for electrolyzers, ranging from $500-1000/kW for large-scale (>20 MW) alkaline or PEM systems in 2024 deployments, with operational expenditures (Opex) dominated by costs at 70-80% of total. At an price of $0.03/kWh and assuming 50-55 kWh/kg H₂ input, levelized costs fall in the $2-3/kg range for high-capacity-factor operation, rising to $4-5/kg with lower utilization or higher Capex; these figures exclude storage costs like compression to 350-700 bar (adding 10-15% penalty) or , which can exceed $1/kg for large volumes. Scaling remains constrained by iridium catalyst scarcity in PEM systems and stack degradation, with real-world efficiencies often 5-10% below lab values due to overpotentials and part-load operation.

Other Specialized Uses

Reversible electrolyzers, which integrate electrolysis and functionalities into a single unit, have been employed in prototyping systems for and conversion. For instance, (PEM) reversible s operate as electrolyzers under applied voltage to produce from and switch to mode for upon reversal of polarity. Solid oxide-based reversible systems, operating at high temperatures, similarly enable electrochemical and power output, with prototypes demonstrating capacities such as 2-kW electrolysis paired with 1-kW performance in alkaline configurations. These setups facilitate testing of bidirectional operation for applications like regenerative s in space or stationary storage. Anodic oxidation via electrolysis serves as a specialized method for treating by oxidizing organic pollutants at the , often generating hydroxyl radicals for degradation without added chemicals. In , boron-doped electrodes achieved near-complete removal of (COD) at natural and ambient temperature, with efficiencies exceeding 90% under galvanostatic conditions. Similar processes have degraded dyes like and methyl blue in synthetic effluents using anodes, attaining over 95% color removal in batch reactors. For , anodic oxidation with reduced organic loads by up to 80% at currents of 1-2 A, targeting recalcitrant pharmaceuticals and pathogens. Electrosynthesis applications include the Kolbe electrolysis, an anodic decarboxylative coupling of carboxylic acids to form hydrocarbons or intermediates, adapted for pharmaceutical precursor production. This process dimerizes carboxylates at platinum anodes, yielding compounds like n-octane from valeric acid with selectivities improved by waveform control to minimize side reactions. In one adaptation, Kolbe decarboxylation enabled synthesis of 2-pyrrolidinone, a key building block for pharmaceutical solvents and drugs, via radical cyclization post-dimerization, offering a greener alternative to traditional routes with yields around 70% under optimized potentials. Such methods leverage electrolysis for precise C-C bond formation in small-scale organic synthesis.

Limitations and Criticisms

Inherent Thermodynamic and Efficiency Constraints

The electrolysis of water requires a minimum Gibbs free energy change of 237 kJ/mol for the production of 1 mol of hydrogen under standard conditions, corresponding to a reversible cell potential of 1.23 V. This sets the theoretical electrical efficiency limit at approximately 83%, calculated as the ratio of ΔG to the higher heating value of hydrogen (286 kJ/mol), assuming no heat recovery. In practice, however, electrode overpotentials—arising from activation barriers at the anode and cathode, mass transport limitations, and ohmic resistances—impose additional voltage requirements, typically elevating the cell voltage to 1.6–2.2 V and reducing efficiencies to 60–80% in low-temperature systems. These overpotentials, combined with from internal resistances, generate irreversible heat losses that cannot be fully harnessed in ambient-temperature electrolyzers operating near room conditions. The excess energy dissipates as , constraining overall cold-endothermic process efficiencies to around 70–75%, as the reaction's endothermic nature (ΔH ≈ 286 kJ/mol) demands external electrical input exceeding the utilizable portion without thermal integration. High-temperature solid oxide electrolysis cells (SOECs), operating at 700–900°C, mitigate some thermodynamic penalties by reducing ΔG through contributions, potentially approaching higher electrical efficiencies by leveraging supplied to offset endothermicity. Nonetheless, the elevated temperatures accelerate material degradation, including cracking, delamination, and , which compromise long-term stack durability and impose inherent operational constraints on sustained performance.

Economic and Scalability Barriers

The levelized cost of (LCOH) produced via electrolysis typically ranges from $3 to $7 per , rendering it uncompetitive with gray derived from reforming, which costs under $2 per as of 2024. constitutes 60-80% of total production costs in electrolytic processes, with variations depending on pricing and grid access, thereby amplifying sensitivity to power market fluctuations. Capital expenditures (Capex) for electrolyzers have shown deflationary trends from 2023 to 2025, driven by scale-up and component optimizations, though absolute levels remain elevated relative to mature technologies. Industry projections target Capex reductions to around $300-500 per kilowatt by mid-decade for alkaline and PEM systems, but these assume aggressive efficiencies not yet realized at gigawatt scales. Operational expenditures are further burdened by stack replacements, occurring every 5-10 years for PEM electrolyzers due to degradation and wear, necessitating recurring investments that elevate long-term costs. Scalability faces material bottlenecks, particularly for PEM electrolyzers reliant on iridium and ruthenium catalysts, where projected demand could reach 32-40 tons of iridium by 2030 amid constrained global supply, exacerbating price volatility and deployment delays. Ruthenium shortages compound these issues in oxygen evolution reaction components, limiting production ramp-up without breakthroughs in low-precious-metal alternatives. Additionally, electrolysis's requirement for continuous high-current operation mismatches intermittent grid supplies from variable renewables, imposing balancing costs that hinder economical large-scale integration without dedicated firm power sources.

Resource and Environmental Realities

Electrolysis for demands substantial , with a stoichiometric minimum of 9 liters of per of H₂ due to the reaction 2H₂O → 2H₂ + O₂, though practical systems consume 10-20 liters per when accounting for cooling, , and inefficiencies in (PEM) or alkaline electrolyzers. In water-scarce arid regions, such as parts of the or , this consumption—equivalent to 9-18 cubic meters of for 1 ton of H₂—exacerbates local shortages, particularly when competing with and municipal needs, and alternatives introduce and energy penalties. When powered by electricity grids with significant fossil fuel contributions, electrolysis inherits substantial carbon emissions, typically 10-20 kg CO₂ equivalent per kg H₂ on average global or regional grids (e.g., U.S. or mixes as of 2023), as the process requires 45-60 kWh per kg H₂ and grid emission factors range from 0.3-0.5 kg CO₂/kWh. This offsets claims of zero-emission unless paired exclusively with low-carbon sources, with coal-dominant grids yielding up to 40 kg CO₂/kg H₂. Even "green" electrolysis reliant on renewables incurs upfront lifecycle emissions from manufacturing solar panels, wind turbines, and electrolyzers—estimated at 5-50 g CO₂/kWh for the electricity pathway—requiring a 1-3 year operational payback period to achieve net-zero status over a 20-30 year lifespan, depending on capacity factors and material sourcing. Renewable infrastructure also expands land footprints, with solar or wind dedicated to electrolysis potentially occupying 10-50 m² per kg H₂ annual capacity (factoring 20-30% capacity factors), limiting scalability in densely populated or ecologically sensitive areas and contributing 10-20% to the total environmental footprint through habitat disruption and biodiversity impacts.

Safety and Reliability Concerns

Electrolysis processes pose significant safety risks primarily from the flammability and potential of gas produced, particularly if it mixes with oxygen or air due to equipment failures such as rupture or gas crossover in the electrolyzer stack. Such mixing can ignite within the cell or during downstream compression, leading to fires or detonations; for instance, a 2007 incident at a power plant involved a during gaseous delivery, highlighting vulnerabilities in handling and storage. Oxygen enrichment in electrolysis cells or exhaust streams exacerbates fire hazards, as elevated oxygen levels above 23.5% increase material reactivity and ignition sensitivity, potentially causing from contaminants like oils or particulates. These risks are often underemphasized compared to dangers, yet they demand stringent purity controls and ventilation to prevent enriched atmospheres in operational environments. Reliability concerns include electrode degradation, notably in proton exchange membrane (PEM) electrolyzers where iridium-based anode catalysts dissolve under anodic conditions, resulting in operational lifetimes typically limited to 20,000–40,000 hours before performance drops necessitate replacement. This dissolution accelerates with dynamic loads or impurities, compromising cell efficiency and requiring frequent maintenance. In seawater electrolysis trials, corrosion from chloride ions and other impurities erodes electrodes, while competitive reactions like hypochlorite formation poison catalysts and reduce selectivity, favoring gas production instead. These issues manifest as rapid material degradation and operational instability, underscoring the challenges in non-purified feedstocks.

Current Research and Developments

Efficiency and Material Advancements

Advancements in design since 2020 have emphasized non-platinum group metal (non-PGM) materials for the reaction (OER) in alkaline electrolysis, particularly NiFe-based compounds, which reduce overpotentials and enhance durability under operational conditions. NiFe (LDHs) and spinels like NiFe₂O₄ have shown superior OER activity, achieving current densities of 320 mA/cm² at 2 V versus (RHE) in 1 M KOH, with sustained performance indicating improved resistance to degradation compared to earlier benchmarks. These catalysts leverage synergies in Ni-Fe electronic structure to lower energy barriers for O-O bond formation, enabling cell voltages below 1.8 V at practical current densities and contributing to overall stack efficiencies exceeding 70% in lab-scale tests. Bipolar stack architectures have scaled electrolysis modules to over 1 MW capacity by optimizing plate materials and flow fields to minimize ohmic losses and gas crossover. These designs interconnect cells in series via conductive bipolar plates, typically or metal composites coated for resistance, supporting uniform current distribution and enabling operation at current densities above 2 A/cm² in (PEM) variants, though alkaline systems often target 0.5-1 A/cm² for longevity. Recent prototypes incorporate advanced sealing and compression techniques, reducing interfacial resistances to below 50 mΩ·cm² and facilitating modular assembly for gigawatt-scale deployment without proportional drops. Anion exchange membrane (AEM) developments, driven by high-throughput materials screening, have yielded membranes with hydroxide conductivities over 100 mS/cm and operational stabilities surpassing 10,000 hours at 60-80°C in alkaline environments. Sustainion®-type AEMs exhibit degradation rates under 1 µV/h at 1.85 V, attributed to quaternary ammonium functionalities resistant to and nucleophilic attack. These membranes, often reinforced with crosslinked polymers, maintain mechanical integrity under differential pressure, enabling AEM electrolyzers to approach PEM performance while using non-PGM electrodes, with cumulative durability data from accelerated stress tests validating projections for multi-year continuous operation.

Alternative Feedstocks and Processes

Direct electrolysis seeks to utilize abundant oceanic as a feedstock for , bypassing costs, yet encounters causal hurdles from ions (≈0.5 M) that thermodynamically and kinetically favor the chlorine evolution reaction (ClER) at potentials overlapping with , diverting current from hydrogen evolution and yielding or Cl₂ byproducts. Precipitates from divalent cations like Mg²⁺ and Ca²⁺ form insulating layers on electrodes during local pH shifts, exacerbating and overpotentials. Empirical pilots in , such as floating platforms, have demonstrated operation but with Faradaic efficiencies for often constrained below 100% due to incomplete ClER suppression, though optimized catalysts achieve >90% in short-term tests; long-term remains limited by these side reactions. Alternative cathodic processes, such as CO₂ electrolysis paired with anodic oxidation, employ CO₂ as a C1 feedstock reduced at cathodes to (CO) or multicarbon products like (C₂H₄), leveraging multi-electron pathways (*CO intermediates dimerize to C-C bonds). However, selectivity suffers from parallel hydrogen evolution and diverse reduction routes (2e⁻ to CO vs. 12e⁻ to ), yielding typical Faradaic efficiencies <50% for under industrially relevant currents (>200 mA/cm²), as confirmed in mechanistic studies of Cu(100) facets and oxide-derived surfaces. Overpotentials exceed 1 V due to CO₂'s low and activation barriers, limiting energy efficiency to ~20-30% without specialized flow cells or promoters. Biomass-assisted hybrid electrolysis substitutes pure water at the with oxygenated organics (e.g., , from waste ), where anodic oxidation potentials (≈0.4-0.8 V vs. RHE) are inherently lower than OER (1.23 V), enabling cell voltage reductions of 0.5-1 V at comparable currents by avoiding O₂ bubble formation and overpotentials. For instance, oxidation on high-OER-potential anodes like Ti/PbO₂ proceeds selectively to quinones or acids, boosting H₂ Faradaic efficiencies >95% while valorizing ; equilibrium shifts can drop theoretical cell voltage to ~0.2 V for certain feedstocks. Challenges include by carbonaceous residues and variable feedstock composition, though pilots confirm savings of 30-50% over standard electrolysis.

Integration with Energy Systems

Electrolysis is frequently coupled with intermittent renewable sources like solar photovoltaic (PV) and wind generation to convert surplus into , addressing overproduction during peak output periods. However, the inherent variability of these sources demands dynamic electrolyzer operation, including frequent ramping and partial loading, which incurs efficiency penalties of approximately 10-20% relative to steady-state baseload conditions due to increased ohmic losses, membrane degradation, and suboptimal stack utilization. Buffering strategies, such as integrating short-term battery storage to smooth input fluctuations, further elevate system costs by 20-50% of electrolyzer while extending response times. Global installed electrolyzer capacity stood at 1.4 GW by the end of 2023, with projections indicating up to 5 GW operational by the end of 2024, primarily driven by projects in and . dominates , achieving an annual electrolyzer production capacity of 39 GW in 2024, supported by scale economies and integration that outpace global demand. In grid contexts, electrolyzers provide balancing services by acting as flexible loads for frequency regulation and peak shaving, potentially absorbing excess renewable output to stabilize networks with high penetration levels. This capability enhances system reliability, as demonstrated in simulations where electrolyzers reduce curtailment by utilizing low-value off-peak power. Yet, hydrogen's indirect climate forcing—via tropospheric chemistry altering lifetimes—yields a 100-year of 12 relative to CO2, meaning leakage rates above 1-2% could negate emissions savings from renewable electrolysis, particularly in unmonitored distribution chains.

Market and Policy Influences

Government subsidies and policy incentives have significantly influenced the deployment of electrolysis for , often distorting market signals by offsetting high production costs. The U.S. of 2022 introduced the 45V Clean Production Tax , offering up to $3 per kilogram for with low lifecycle emissions, which has spurred announcements of new projects but relies on ongoing fiscal support to achieve viability. Without such subsidies, costs range from $3.50 to $8 per kilogram, far exceeding gray hydrogen at $1.50 to $6.40 per kilogram, rendering electrolysis uncompetitive absent access to below $0.03 per kWh. Announced electrolyzer manufacturing capacity has expanded rapidly, with global installations reaching 1.4 GW by end-2023 and projections for annual additions climbing toward 95 GW by 2030, yet operational low-emissions remains below 1% of total production as of 2024. Implementation gaps persist, with only 7% of announced capacity realized by 2023, highlighting over-optimism in policy-driven forecasts. Electrolytic is projected to constitute less than 4% of global supply by 2030, constrained by scalability and the dominance of fossil-based methods. Policy emphasis on "green" hydrogen overlooks its frequent grid dependency, where electricity sourced from non-renewable grids undermines emissions reductions and prolongs fossil fuel reliance in power generation. This approach risks delaying the phase-out of unabated , as subsidized projects prioritize volume over genuine decarbonization. For most applications, direct offers superior , avoiding electrolysis losses of 25-40% in conversion and storage, making it preferable where feasible over intermediaries. Market-driven adoption would likely favor such alternatives unless electrolysis achieves cost parity through technological breakthroughs independent of mandates.

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