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Electrolysis of water

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This article is about the electrochemical system. For related systems, see Photoelectrolysis of water and Photocatalytic water splitting.

Simple setup for demonstration of electrolysis of water at home
An AA battery in a glass of tap water with salt showing hydrogen produced at the negative terminal

Electrolysis of water is using electricity to split water into oxygen (O
2) and hydrogen (H
2) gas by electrolysis. Hydrogen gas released in this way can be used as hydrogen fuel, but must be kept apart from the oxygen as the mixture would be extremely explosive. Separately pressurised into convenient 'tanks' or 'gas bottles', hydrogen can be used for oxyhydrogen welding and other applications, as the hydrogen / oxygen flame can reach approximately 2,800°C.

Water electrolysis requires a minimum potential difference of 1.23 volts, although at that voltage external heat is also required. Typically 1.5 volts is required. Electrolysis is rare in industrial applications since hydrogen can be produced less expensively from fossil fuels.[1] Most of the time, hydrogen is made by splitting methane (CH4) into carbon dioxide (CO2) and hydrogen (H2). It’s a carbon-intensive process that means for every kilogram of “grey” hydrogen produced, approximately 10 kilograms of CO2 are emitted into the atmosphere.[2]

History

[edit]
Device invented by Johann Wilhelm Ritter to develop the electrolysis of water

In 1789, Jan Rudolph Deiman and Adriaan Paets van Troostwijk used an electrostatic machine to make electricity that was discharged on gold electrodes in a Leyden jar.[3] In 1800, Alessandro Volta invented the voltaic pile, while a few weeks later English scientists William Nicholson and Anthony Carlisle used it to electrolyse water. In 1806 Humphry Davy reported the results of extensive distilled water electrolysis experiments, concluding that nitric acid was produced at the anode from dissolved atmospheric nitrogen. He used a high voltage battery and non-reactive electrodes and vessels such as gold electrode cones that doubled as vessels bridged by damp asbestos.[4] Zénobe Gramme invented the Gramme machine in 1869, making electrolysis a cheap method for hydrogen production. A method of industrial synthesis of hydrogen and oxygen through electrolysis was developed by Dmitry Lachinov in 1888.[5]

Principles

[edit]

A DC electrical power source is connected to two electrodes, or two plates (typically made from an inert metal such as platinum or iridium) that are placed in the water. Hydrogen appears at the cathode (where electrons enter the water), and oxygen at the anode.[6] Assuming ideal faradaic efficiency, the amount of hydrogen generated is twice the amount of oxygen, and both are proportional to the total electrical charge conducted by the solution.[7] However, in many cells competing side reactions occur, resulting in additional products and less than ideal faradaic efficiency.[citation needed]

Electrolysis of pure water requires excess energy in the form of overpotential to overcome various activation barriers. Without the excess energy, electrolysis occurs slowly or not at all. This is in part due to the limited self-ionization of water.

Pure water has an electrical conductivity about one hundred thousandth that of seawater.[8][9][10]

Efficiency is increased through the addition of an electrolyte (such as a salt, an acid or a base) and electrocatalysts.

Equations

[edit]
Diagram showing the overall chemical equation.

In pure water at the negatively charged cathode, a reduction reaction takes place, with electrons (e−) from the cathode being given to hydrogen cations to form hydrogen gas. At the positively charged anode, an oxidation reaction occurs, generating oxygen gas and giving electrons to the anode to complete the circuit.

The two half-reactions, reduction and oxidation, are coupled to form a balanced system. In order to balance each half-reaction, the water needs to be acidic or basic. In the presence of acid, the equations are:

Cathode (reduction): 2H+(aq) + 2e− → H2(g)
Anode (oxidation): H2O(l) → 1/2 O2(g) + 2 H+(aq) + 2e−

In the presence of base, the equations are:

Cathode (reduction): 2 H2O(l) + 2e− → H2(g) + 2 OH−(aq)
Anode (oxidation): 2 OH−(aq) → 1/2 O2(g) + H2O(l) + 2 e−

Combining either half reaction pair yields the same overall decomposition of water into oxygen and hydrogen:

2 H2O(l) → 2 H2(g) + O2(g)

The number of hydrogen molecules produced is thus twice the number of oxygen molecules, in keeping with the facts that both hydrogen and oxygen are diatomic molecules and water molecules contain twice as many hydrogen atoms as oxygen atoms. Assuming equal temperature and pressure for both gases, volume is proportional to moles, so twice as large a volume of hydrogen gas is produced as oxygen gas. The number of electrons pushed through the water is twice the number of generated hydrogen molecules and four times the number of generated oxygen molecules.

Thermodynamics

[edit]
Pourbaix diagram for water, including equilibrium regions for water, oxygen and hydrogen at STP. The vertical scale is the electrode potential of hydrogen or non-interacting electrode relative to an SHE electrode, the horizontal scale is the pH of the electrolyte (otherwise non-interacting). Neglecting overpotential, above the top line the equilibrium condition is oxygen gas, and oxygen will bubble off of the electrode until equilibrium is reached. Likewise, below the bottom line, the equilibrium condition is hydrogen gas, and hydrogen will bubble off of the electrode until equilibrium is reached.

The decomposition of pure water into hydrogen and oxygen at standard temperature and pressure is not favorable in thermodynamic terms.

Anode (oxidation): 2 H2O(l) → O2(g) + 4 H+(aq) + 4e−    Eo = +1.23 V (for the reduction half-equation)
Cathode (reduction): 2 H+(aq) + 2e− → H2(g) Eo = 0.00 V

Thus, the standard potential of the water electrolysis cell (Eocell = Eocathode − Eoanode) is −1.229 V at 25 °C at pH 0 ([H+] = 1.0 M). At 25 °C with pH 7 ([H+] = 1.0×10−7 M), the potential is unchanged based on the Nernst equation. The thermodynamic standard cell potential can be obtained from standard-state free energy calculations to find ΔG° and then using the equation: ΔG°= −n F E° (where E° is the cell potential and F the Faraday constant, 96,485 C/mol). For two water molecules electrolysed and hence two hydrogen molecules formed, n = 4, and

  • ΔG° = 474.48 kJ/2 mol(water) = 237.24 kJ/mol(water)
  • ΔS° = 163 J/K mol(water)
  • ΔH° = 571.66 kJ/2 mol(water) = 285.83 kJ/mol(water)
  • and 141.86 kJ/g(H2).

However, calculations regarding individual electrode equilibrium potentials requires corrections to account for the activity coefficients.[11] In practice when an electrochemical cell is "driven" toward completion by applying reasonable potential, it is kinetically controlled. Therefore, activation energy, ion mobility (diffusion) and concentration, wire resistance, surface hindrance including bubble formation (blocks electrode area), and entropy, require greater potential to overcome. The amount of increase in required potential is termed the overpotential.

Electrolyte

[edit]
Hofmann voltameter connected to a direct current power supply

Electrolysis in pure water consumes/reduces H+ cations at the cathode and consumes/oxidizes hydroxide (OH−) anions at the anode. This can be verified by adding a pH indicator to the water: Water near the cathode is basic while water near the anode is acidic. The hydroxides OH− that approach the anode mostly combine with the positive hydronium ions (H3O+) to form water. The positive hydronium ions that approach the cathode mostly combine with negative hydroxide ions to form water. Relatively few hydroniums/hydroxide ions reach the cathode/anode. This can cause overpotential at both electrodes.

Pure water has a charge carrier density similar to semiconductors[12][page needed] since it has a low autoionization, Kw = 1.0×10−14 at room temperature and thus pure water conducts current poorly, 0.055 μS/cm.[13] Unless a large potential is applied to increase the autoionization of water, electrolysis of pure water proceeds slowly, limited by the overall conductivity.

An aqueous electrolyte can considerably raise conductivity. The electrolyte disassociates into cations and anions; the anions rush towards the anode and neutralize the buildup of positively charged H+ there; similarly, the cations rush towards the cathode and neutralize the buildup of negatively charged OH− there. This allows the continuous flow of electricity.[14]

Anions from the electrolyte compete with the hydroxide ions to give up an electron. An electrolyte anion with less standard electrode potential than hydroxide will be oxidized instead of the hydroxide, producing no oxygen gas. Likewise, a cation with a greater standard electrode potential than a hydrogen ion will be reduced instead of hydrogen.

Various cations have lower electrode potential than H+ and are therefore suitable for use as electrolyte cations: Li+, Rb+, K+, Cs+, Ba2+, Sr2+, Ca2+, Na+, and Mg2+. Sodium and potassium are common choices,[15] as they form inexpensive, soluble salts.

If an acid is used as the electrolyte, the cation is H+, and no competitor for the H+ is created by disassociating water. The most commonly used anion is sulfate (SO2−
4), as it is difficult to oxidize. The standard potential for oxidation of this ion to the peroxydisulfate ion is +2.010 volts.[16]

Strong acids such as sulfuric acid (H2SO4), and strong bases such as potassium hydroxide (KOH), and sodium hydroxide (NaOH) are common choices as electrolytes due to their strong conducting abilities.

A solid polymer electrolyte can be used such as Nafion and when applied with an appropriate catalyst on each side of the membrane can efficiently electrolyze with as little as 1.5  volts. Several commercial electrolysis systems use solid electrolytes.[17]

Pure water

[edit]
Main article: Virtual breakdown mechanism

Electrolyte-free pure water electrolysis has been achieved via deep-sub-Debye-length nanogap electrochemical cells. When the gap between cathode and anode are smaller than Debye-length (1 micron in pure water, around 220 nm in distilled water), the double layer regions from two electrodes can overlap, leading to a uniformly high electric field distributed across the entire gap. Such a high electric field can significantly enhance ion transport (mainly due to migration), further enhancing self-ionization, continuing the reaction and showing little resistance between the two electrodes. In this case, the two half-reactions are coupled and limited by electron-transfer steps (the electrolysis current is saturated at shorter electrode distances).[18]

Seawater

[edit]

Ambient seawater presents challenges because of the presence of salt and other impurities. Approaches may or may not involve desalination before electrolysis. Traditional electrolysis produces toxic and corrosive chlorine ions (e.g., Cl−
 and ClO−
).[19][20] Multiple methods have been advanced for electrolysing unprocessed seawater. Typical proton exchange membrane (PEM) electrolysers require desalination.

Indirect seawater electrolysis involves two steps: desalting seawater using a pre-treatment device and then producing hydrogen through traditional water electrolysis. This method improves efficiency, reduces corrosion, and extends catalyst lifespan.[21] Some argue that the costs of seawater desalination are relatively small compared to water splitting, suggesting that research should focus on developing more efficient two-step desalination-coupled water splitting processes.[22][23]

However, indirect seawater electrolysis plants require more space, energy, and more maintenance, and some believe that the water purity achieved through seawater reverse osmosis (SWRO) may not be sufficient, necessitating additional equipment and cost.[21] In contrast, direct seawater electrolysis skips the pre-treatment step and introduces seawater directly into the electrolyzer to produce hydrogen. This approach is seen as more promising due to limited freshwater resources, the need to prioritize basic human needs, and the potential to reduce energy consumption and costs.[23][24][25] Membranes are critical for the efficiency of electrolysis, but they can be negatively affected by foreign ions in seawater, shortening their lifespan and reducing the efficiency of the electrolysis process.[26]

One approach involves combining forward osmosis membranes with water splitting to produce hydrogen continuously from impure water sources. Water splitting generates a concentration gradient balanced by water influx via forward osmosis, allowing for continual extraction of pure water. However, this configuration has challenges such as the potential for Cl ions to pass through the membrane and cause damage, as well as the risk of hydrogen and oxygen mixing without a separator.[27]

To address these issues, a low-cost semipermeable membrane was introduced between the electrodes to separate the generated gases, reducing membrane costs and minimizing Cl oxidation. Additionally, research shows that using transition metal-based materials can support water electrolysis efficiently.[28] Some studies have explored the use of low-cost reverse osmosis membranes (<10$/m2) to replace expensive ion exchange membranes (500-1000$/m2). The use of reverse osmosis membranes becomes economically attractive in water electrolyzer systems as opposed to ion exchange membranes due to their cost-effectiveness and the high proton selectivity they offer for cation salts, especially when high-concentration electrolytes are employed. [29]

An alternative method involves employing a hydrophobic membrane to prevent ions from entering the cell stack. This method combines a hydrophobic porous polytetrafluoroethylene (PTFE) waterproof breathable membrane with a self-dampening electrolyte, utilizing a hygroscopic sulfuric acid solution with a commercial alkaline electrolyzer to generate hydrogen gas from seawater. At a larger scale, this seawater electrolysis system can consistently produce 386 L of H2 per hour for over 3200 hours without experiencing significant catalyst corrosion or membrane wetting. The process capitalizes on the disparity in water vapor pressure between seawater and the self-dampening electrolyte to drive seawater evaporation and water vapor diffusion, followed by the liquefaction of the adsorbed water vapor on the self-dampening electrolyte.[30][21]

Techniques

[edit]

As of 2022, commercial electrolysis requires around 53 kWh of electricity to produce one kg of hydrogen, which holds 39.4 kWh (HHV) of energy.[31]

Fundamental demonstration

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Two leads, running from the terminals of a battery, placed in a cup of water with a quantity of electrolyte establish conductivity. Using NaCl (salt) in an electrolyte solution yields chlorine gas rather than oxygen due to a competing half-reaction. Sodium bicarbonate (baking soda) instead yields hydrogen, and carbon dioxide for as long as the bicarbonate anion stays in solution.

Match test used to detect the presence of hydrogen gas

Hofmann voltameter

[edit]
Main article: Hofmann voltameter

The Hofmann voltameter is a small-scale electrolytic cell. It consists of three joined upright cylinders. The inner cylinder is open at the top to allow the addition of water and electrolyte. A platinum electrode (plate or honeycomb) is placed at the bottom of each of the two side cylinders, connected to the terminals of an electricity source. The generated gases displace water and collect at the top of the two outer tubes, where it can be drawn off with a stopcock.

High-pressure

[edit]
Main article: High-pressure electrolysis

High-pressure electrolysis involves compressed hydrogen output around 12–20 MPa (120–200 Bar, 1740–2900 psi).[32] By pressurising the hydrogen in the electrolyser, the need for an external hydrogen compressor is eliminated. The average energy consumption is around 3%.[33]

High-temperature

[edit]
Main article: High-temperature electrolysis
Theoretical thermal water splitting efficiencies.[34]
60% efficient at 1000°C
Steam reforming of hydrocarbons to hydrogen is 70-85% efficient[35]

High-temperature electrolysis (also HTE or steam electrolysis) is more efficient at higher temperatures. A heat engine supplies some of the energy, which is typically cheaper than electricity [36][37]

Alkaline water electrolysis

[edit]
This section is an excerpt from Alkaline water electrolysis.[edit]

Alkaline water electrolysis is a type of electrolysis that is characterized by having two electrodes operating in a liquid alkaline electrolyte. Commonly, a solution of potassium hydroxide (KOH) or sodium hydroxide (NaOH) at 25-40 wt% is used.[38] These electrodes are separated by a diaphragm, separating the product gases and transporting the hydroxide ions (OH−) from one electrode to the other.[39][40] A recent comparison showed that state-of-the-art nickel based water electrolysers with alkaline electrolytes lead to competitive or even better efficiencies than acidic polymer electrolyte membrane water electrolysis with platinum group metal based electrocatalysts.[41]

The technology has a long history in the chemical industry. The first large-scale demand for hydrogen emerged in late 19th century for lighter-than-air aircraft, and before the advent of steam reforming in the 1930s, the technique was competitive.[citation needed]

Hydrogen-based technologies have evolved significantly since the initial discovery of hydrogen and its early application as a buoyant gas approximately 250 years ago. In 1804, the Swiss inventor Francois Isaac de Rivaz secured a patent for the inaugural hydrogen-powered vehicle. This prototype, equipped with a four-wheel design, utilised an internal combustion engine (ICE) fuelled by a mixture of hydrogen and oxygen gases. The hydrogen fuel was stored in a balloon, and ignition was achieved through an electrical starter known as a Volta starter. The combustion process propelled the piston within the cylinder, which, upon descending, activated a wheel through a ratchet mechanism. This invention could be viewed as an early embodiment of a system comprising hydrogen storage, conduits, valves, and a conversion device.[42]

Approximately four decades after the military scientist Ritter developed the first electrolyser, the chemists Schoenbein and Sir Grove independently identified and showcased the fuel cell concept. This technology operates in reverse to electrolysis around the year 1839. This discovery marked a significant milestone in the field of hydrogen technology, demonstrating the potential for hydrogen as a source of clean energy.[42]

Proton exchange membrane

[edit]
Main article: Proton exchange membrane electrolysis

A proton-exchange membrane electrolyser separates reactants and transports protons while blocking a direct electronic pathway through the membrane. PEM fuel cells use a solid polymer membrane (a thin plastic film) which is permeable to hydrogen ions (protons) when it is saturated with water, but does not conduct electrons.

It uses a proton-exchange membrane, or polymer-electrolyte membrane (PEM), which is a semipermeable membrane generally made from ionomers and designed to conduct protons while acting as an insulator and reactant barrier, e.g. to oxygen and hydrogen gas.[43] PEM fuel cells use a solid polymer membrane (a thin plastic film) that is permeable to protons when saturated with water, but does not conduct electrons. Proton-exchange membranes are primarily characterized by proton conductivity (σ), methanol permeability (P), and thermal stability.[44]

PEMs can be made from either pure polymer or from composite membranes, where other materials are embedded in a polymer matrix. One of the most common commercially available materials is the fluoropolymer (PFSA)[45] Nafion.[46] Nafion is an ionomer with a perfluorinated backbone such as Teflon.[47] Many other structural motifs are used to make ionomers for proton-exchange membranes. Many use polyaromatic polymers, while others use partially fluorinated polymers.

Anion exchange membrane

[edit]
Main article: Anion exchange membrane electrolysis

Anion exchange membrane electrolysis employs an anion-exchange membrane (AEM) to achieve the separation of products, provide electrical insulation between electrodes, and facilitate ion conduction. In contrast to PEM electrolysis, AEM electrolysis allows for the conduction of hydroxide ions. A noteworthy benefit of AEM-based water electrolysis is the elimination of the need for expensive noble metal catalysts, as cost-effective transition metal catalysts can be utilized in their place.[48][49]

Supercritical water

[edit]

Supercritical water electrolysis (SWE) uses water in a supercritical state. Supercritical water requires less energy, therefore reducing costs. It operates at >375 °C, which reduces thermodynamic barriers and increases kinetics, improving ionic conductivity over liquid or gaseous water, which reduces ohmic losses. Benefits include improved electrical efficiency, >221bar pressurised delivery of product gases, ability to operate at high current densities and low dependence on precious metal catalysts. As of 2021 commercial SWE equipment was not available.[50]

Nickel/iron

[edit]

In 2014, researchers announced electrolysis using nickel and iron catalysts rather than precious metals. Nickel-metal/nickel-oxide structure is more active than nickel metal or nickel oxide alone. The catalyst significantly lowers the required voltage.[51][52] Nickel–iron batteries are under investigation for use as combined batteries and electrolysers. Those "battolysers" could be charged and discharged like conventional batteries, and would produce hydrogen when fully charged.[53]

In 2023, researchers in Australia announced the use of a porous sheet of nitrogen-doped nickel molybdenum phosphide catalyst. The nitrogen doping increases conductivity and optimizes electronic density and surface chemistry. This produces additional catalytic sites. The nitrogen bonds to the surface metals and has electro-negative properties that help exclude unwanted ions and molecules, while phosphate, sulfate, nitrate and hydroxyl surface ions block chlorine and prevent corrosion. 10 mA/cm2 can be achieved using 1.52 and 1.55 V in alkaline electrolyte and seawater, respectively.[54]

Nanogap electrochemical cells

[edit]
Main article: Virtual breakdown mechanism

In 2017, researchers reported nanogap electrochemical cells that achieved high-efficiency electrolyte-free pure water electrolysis at ambient temperature. In these cells, the two electrodes are so close to each other (smaller than Debye-length) that the mass transport rate can be higher than the electron-transfer rate, leading to two half-reactions coupled together and limited by the electron-transfer step. Experiments show that the electrical current density can be larger than that from 1  mol/L sodium hydroxide solution. Its "Virtual Breakdown Mechanism", is completely different from traditional electrochemical theory, due to such nanogap size effects.[18]

Capillary fed

[edit]

A capillary-fed electrolyzer cell is claimed to require only 41.5 kWh to produce 1 kg of hydrogen. The water electrolyte is isolated from the electrodes by a porous, hydrophilic separator. The water is drawn into the electrolyzer by capillary action, while the electrolyzed gases pass out on either side. It extends PEM technology by eliminating bubbles that reduce the contact between the electrodes and the electrolyte, reducing efficiency. The design is claimed to operate at 98% energy efficiency (higher heating value of hydrogen). The design forgoes water circulation, separator tanks, and other mechanism and can be air- or radiatively cooled.[31][55] The effect of the build-up of impurities in the cell from those initially present in the feed water is not yet available.

Applications

[edit]

About five percent of hydrogen gas produced worldwide is created by electrolysis. The vast majority of current industrial hydrogen production is from natural gas in the steam reforming process, or from the partial oxidation of coal or heavy hydrocarbons. The majority[citation needed] of the hydrogen produced through electrolysis is a side product in the production of chlorine and caustic soda. This is a prime example of a competing for side reaction.

2NaCl + 2H2O → Cl2 + H2 + 2NaOH

In the chloralkali process (electrolysis of brine) a water/sodium chloride mixture is only half the electrolysis of water since the chloride ions are oxidized to chlorine rather than water being oxidized to oxygen. Thermodynamically, this would not be expected since the oxidation potential of the chloride ion is less than that of water, but the rate of the chloride reaction is much greater than that of water, causing it to predominate. The hydrogen produced from this process is either burned (converting it back to water), used for the production of specialty chemicals, or various other small-scale applications.

Water electrolysis is also used to generate oxygen for the International Space Station.[56][57]

Many industrial electrolysis cells are similar to Hofmann voltameters, with platinum plates or honeycombs as electrodes. Generally, hydrogen is produced for point of use applications such as oxyhydrogen torches or when high purity hydrogen or oxygen is desired. The vast majority of hydrogen is produced from hydrocarbons and as a result, contains trace amounts of carbon monoxide among other impurities. The carbon monoxide impurity can be detrimental to various systems including many fuel cells.

As electrolysers can be ramped down they might in future be used to cope with electricity supply demand mismatch.[58]

Efficiency

[edit]

Industrial output

[edit]
Illustrating inputs and outputs of simple electrolysis of water, for production of hydrogen.

Efficiency of modern hydrogen generators is measured by energy consumed per standard volume of hydrogen (MJ/m3), assuming standard temperature and pressure of the H2. The lower the energy used by a generator, the higher its efficiency would be; a 100%-efficient electrolyser would consume 39.4 kilowatt-hours per kilogram (142 MJ/kg) (higher heating value) of hydrogen,[59] 12,749 joules per litre (12.75 MJ/m3). Practical electrolysis (using a rotating electrolyser at 15 bar pressure) may consume 50 kW⋅h/kg (180 MJ/kg), and a further 15 kW⋅h (54 MJ) if the hydrogen is compressed for use in hydrogen cars.[60] By adding external heat at 150 °C (302 °F), electricity consumption may be reduced.[61]

There are three main technologies available on the market: alkaline, proton exchange membrane (PEM), and solid oxide electrolyzers.

Alkaline electrolyzers are cheaper in terms of investment (they generally use nickel catalysts), but least efficient. PEM electrolyzers are more expensive (they generally use expensive platinum-group metal catalysts) but are more efficient and can operate at higher current densities, and can, therefore, be possibly cheaper if the hydrogen production is large enough. Solid oxide electrolyzer cells (SOEC) are the third most common type of electrolysis, and the most expensive, and use high operating temperatures to increase efficiency. The theoretical electrical efficiency of SOEC is close to 100% at 90% hydrogen production.[62] Degradation of the system over time does not affect the efficiency of SOEC electrolyzers initially unlike PEM and alkaline electrolyzers. As the SOEC system degrades, the cell voltage increases, producing more heat in the system naturally. Due to this, less energy is required to keep the system hot, which will make up for the energy losses from dramatic degradation initially.[63] SOEC requires replacement of the stack after some years of degradation.

Efficiency

[edit]

Electrolyzer vendors provide efficiencies based on enthalpy. To assess the claimed efficiency of an electrolyzer it is important to establish how it was defined by the vendor (i.e. what enthalpy value, what current density, etc.).

Conventional alkaline electrolysis has an efficiency of about 70%.[64] Accounting for the accepted use of the higher heating value (because inefficiency via heat can be redirected back into the system to create the steam required by the catalyst), average working efficiencies for PEM electrolysis are around 80%.[65][66] This is expected to increase to between 82 and 86%[67] before 2030. Theoretical efficiency for PEM electrolysers are predicted up to 94%.[68]

In 2024, Australian company Hysata announced a device capable of 95% efficiency relative to the higher heating value of hydrogen. Conventional systems consume 52.5 kWh to produce hydrogen that can store 39.4 kWh of energy (1 kg). Its technology requires only 41.5 kWh to produce 1 kg. It uses a capillary-fed electrolyzer to eliminate hydrogen and oxygen bubbles in the fluid electrolyte. Bubbles are non-conductive, and can stick to electrodes, reducing electrode exposure to the electrolyte, increasing resistance. Hysata places the electrolyte at the bottom of the device. Capillary action draws it through a porous, hydrophilic separator between the electrodes. Each electrode has complete contact with the electrolyte on the inner side, and a dry chamber on the outer side.[69][70] The effect of the build-up of impurities in the cell from those initially present in the feed water is not yet available.

Cost

[edit]

Calculating cost is complicated,[71] and a market price barely exists.[72] Considering the industrial production of hydrogen, and using current best processes for water electrolysis (PEM or alkaline electrolysis) which have an effective electrical efficiency of 70–80%,[68][73][74] producing 1 kg of hydrogen (which has a specific energy of 143 MJ/kg) requires 50–55 kW⋅h (180–200 MJ) of electricity. At an electricity cost of $0.06/kW·h, as set out in the US Department of Energy hydrogen production targets for 2015,[75] the hydrogen cost is $3/kg. Equipment cost depends on mass production. Operating cost depends on electricity cost for about half of the levelised product price.[72][71]

H2 production cost ($-gge untaxed) at varying natural gas prices

Comparison with steam-methane-reformed (SMR) hydrogen

[edit]

With the range of natural gas prices from 2016 as shown in the graph (Hydrogen Production Tech Team Roadmap, November 2017) putting the cost of steam-methane-reformed (SMR) hydrogen at between $1.20 and $1.50, the cost price of hydrogen via electrolysis is still over double 2015 DOE hydrogen target prices. The US DOE target price for hydrogen in 2020 is $2.30/kg, requiring an electricity cost of $0.037/kW·h, which is achievable given 2018 PPA tenders[76] for wind and solar in many regions. This puts the $4/gasoline gallon equivalent (gge) H2 dispensed objective well within reach, and close to a slightly elevated natural gas production cost for SMR.

In other parts of the world, the price of SMR hydrogen is between $1–3/kg on average. This makes production of hydrogen via electrolysis cost competitive in many regions already, as outlined by Nel Hydrogen[77] and others, including an article by the IEA[78] examining the conditions which could lead to a competitive advantage for electrolysis. The large price increase of gas during the 2021–2022 global energy crisis made hydrogen electrolysis economic in some parts of the world.[79]

Facilities

[edit]

Some large industrial electrolyzers are operating at several megawatts. As of 2022, the largest is a 150 MW alkaline facility in Ningxia, China, with a capacity up to 23,000 tonnes per year.[80] While higher-efficiency Western electrolysis equipment can cost $1,200/kW, lower-efficiency Chinese equipment can cost $300/kW, but with a lower lifetime of 60,000 hours.[81]

As of 2022, different analysts predict annual manufacture of equipment by 2030 as 47 GW, 104 GW and 180 GW, respectively.[82]

Overpotential

[edit]

Real water electrolyzers require higher voltages for the reaction to proceed. The part that exceeds 1.23 V[83] is called overpotential or overvoltage, and represents any kind of loss and nonideality in the electrochemical process.

For a well designed cell the largest overpotential is the reaction overpotential for the four-electron oxidation of water to oxygen at the anode; electrocatalysts can facilitate this reaction, and platinum alloys are the state of the art for this oxidation. Developing a cheap, effective electrocatalyst for this reaction would be a great advance, and is a topic of current research; there are many approaches, among them a 30-year-old recipe for molybdenum sulfide,[84] graphene quantum dots,[85] carbon nanotubes,[52] perovskite,[86] and nickel/nickel-oxide.[87][88] Trimolybdenum phosphide (Mo3P) has been recently found as a promising nonprecious metal and earth‐abundant candidate with outstanding catalytic properties that can be used for electrocatalytic processes. The catalytic performance of Mo3P nanoparticles is tested in the hydrogen evolution reaction (HER), indicating an onset potential of as low as 21 mV, H2 formation rate, and exchange current density of 214.7 μmol/(s·g) cat (at only 100 mV overpotential) and 279.07 μA/cm2, respectively, which are among the closest values yet observed to platinum.[89][90] The simpler two-electron reaction to produce hydrogen at the cathode can be electrocatalyzed with almost no overpotential by platinum, or in theory a hydrogenase enzyme. If other, less effective, materials are used for the cathode (e.g. graphite), large overpotentials will appear.

Thermodynamics

[edit]

The electrolysis of water in standard conditions requires a theoretical minimum of 237 kJ of electrical energy input to dissociate each mole of water, which is the standard Gibbs free energy of formation of water. It also requires thermal energy to balance the change in entropy of the reaction. Therefore, the process cannot proceed at constant temperature at electrical energy inputs below 286 kJ per mol if no external thermal energy is added.

Since each mole of water requires two moles of electrons, and given that the Faraday constant F represents the charge of a mole of electrons (96485 C/mol), it follows that the minimum voltage necessary for electrolysis is about 1.23 V.[91] If electrolysis is carried out at high temperature, this voltage reduces. This effectively allows the electrolyser to operate at more than 100% electrical efficiency. In electrochemical systems this means that heat must be supplied to the reactor to sustain the reaction. In this way thermal energy can be used for part of the electrolysis energy requirement.[92] In a similar way the required voltage can be reduced (below 1 V) if fuels (such as carbon, alcohol, biomass) are reacted with water (PEM based electrolyzer in low temperature) or oxygen ions (solid oxide electrolyte based electrolyzer in high temperature). This results in some of the fuel's energy being used to "assist" the electrolysis process and can reduce the overall cost of hydrogen produced.[93]

However, observing the entropy component (and other losses), voltages over 1.48 V are required for the reaction to proceed at practical current densities (the thermoneutral voltage).

In the case of water electrolysis, Gibbs free energy represents the minimum work necessary for the reaction to proceed, and the reaction enthalpy is the amount of energy (both work and heat) that has to be provided so the reaction products are at the same temperature as the reactant (i.e. standard temperature for the values given above). Potentially, an electrolyzer operating at 1.48 V would operate isothermally at a temperature of 25°C as the electrical energy supplied would be equal to the enthalpy (heat) of water decomposition and this would require 20% more electrical energy than the minimum.

See also

[edit]

References

[edit]
  1. ^ "Hydrogen Basics — Production". Florida Solar Energy Center. 2007. Archived from the original on 18 February 2008. Retrieved 5 February 2008.
  2. ^ Cuff, Madeleine (28 August 2024). "Is ultra cheap green hydrogen on the horizon?". New Scientist. Retrieved 29 August 2024.
  3. ^ Levie, R. de (October 1999). "The electrolysis of water". Journal of Electroanalytical Chemistry. 476 (1): 92–93. doi:10.1016/S0022-0728(99)00365-4.
  4. ^ Davy, John, ed. (1839). "On Some Chemical Agencies of Electricity". The Collected Works of Sir Humphry Davy. Vol. 5. pp. 1–12.
  5. ^ "Lachinov Dmitry Aleksandrovich". Great Cyrill and Methodius Encyclopedia (in Russian). Archived from the original on 26 July 2011.
  6. ^ Zumdahl, Steven S.; Zumdahl, Susan A. (1 January 2013). Chemistry (9th ed.). Cengage Learning. p. 30. ISBN 978-1-13-361109-7.
  7. ^ Carmo, M; Fritz D; Mergel J; Stolten D (2013). "A comprehensive review on PEM water electrolysis". Journal of Hydrogen Energy. 38 (12): 4901–4934. doi:10.1016/j.ijhydene.2013.01.151.
  8. ^ "5.9 Conductivity | Monitoring & Assessment | US EPA". United States Environmental Protection Agency Web Archive. Archived from the original on 30 July 2024. Retrieved 20 October 2024.
  9. ^ Unknown. "Water quality standards" (PDF). Mary River Catchment Coordinating Committee. Archived (PDF) from the original on 3 October 2024. Retrieved 20 October 2024.
  10. ^ Clean Water Team (CWT) 2004. Electrical conductivity/salinity Fact Sheet, FS3.1.3.0(EC). in: The Clean Water Team Guidance Compendium for Watershed Monitoring and Assessment, Version 2.0. Division of Water Quality, California State Water Resources Control Board (SWRCB), Sacramento, CA.
  11. ^ Colli, A.N.; et al. (2019). "Non-Precious Electrodes for Practical Alkaline Water Electrolysis". Materials. 12 (8): 1336. Bibcode:2019Mate...12.1336C. doi:10.3390/ma12081336. PMC 6515460. PMID 31022944.
  12. ^ Fuller, C. S. (1959). "Defect Interactions in Semiconductors". In Hannay, N. B. (ed.). Semiconductors. New York: Reinhold. pp. 192–221.
  13. ^ Light, Truman S.; Licht, Stuart; Bevilacqua, Anthony C.; Morash, Kenneth R. (1 January 2005). "The Fundamental Conductivity and Resistivity of Water". Electrochemical and Solid-State Letters. 8 (1): E16–E19. doi:10.1149/1.1836121. ISSN 1099-0062. S2CID 54511887.
  14. ^ PAULING, LINUS (1953). "Section 15-2". General Chemistry (2nd ed.).
  15. ^ Chatenet, Marian; Pollet, Bruno G.; Dekel, Dario R.; Dionigi, Fabio; Deseure, Jonathan; Millet, Pierre; Braatz, Richard D.; Bazant, Martin Z.; Eikerling, Michael; Staffell, Iain; Balcombe, Paul; Shao-Horn, Yang; Schäfer, Helmut (2022). "Water electrolysis: from textbook knowledge to the latest scientific strategies and industrial developments". Chemical Society Reviews. 51 (11): 4583–4762. doi:10.1039/d0cs01079k. PMC 9332215. PMID 35575644.
  16. ^ Haynes, William M. (2012). CRC handbook of chemistry and physics : a ready-reference book of chemical and physical data (93rd, 2012-2013 ed.). Boca Raton, Fla.: CRC. ISBN 9781439880494. OCLC 793213751.
  17. ^ Badwal, SPS; Giddey S; Munnings C (2012). "Hydrogen production via solid electrolytic routes". WIREs Energy and Environment. 2 (5): 473–487. Bibcode:2013WIREE...2..473B. doi:10.1002/wene.50. S2CID 135539661. Archived from the original on 2 June 2013. Retrieved 23 January 2013.
  18. ^ a b Wang, Yifei; Narayanan, S. R.; Wu, Wei (11 July 2017). "Field-Assisted Splitting of Pure Water Based on Deep-Sub-Debye-Length Nanogap Electrochemical Cells". ACS Nano. 11 (8): 8421–8428. doi:10.1021/acsnano.7b04038. ISSN 1936-0851. PMID 28686412.
  19. ^ Kuang, Yun; Kenney, Michael J.; Meng, Yongtao; Hung, Wei-Hsuan; Liu, Yijin; Huang, Jianan Erick; Prasanna, Rohit; Li, Pengsong; Li, Yaping; Wang, Lei; Lin, Meng-Chang; McGehee, Michael D.; Sun, Xiaoming; Dai, Hongjie (2 April 2019). "Solar-driven, highly sustained splitting of seawater into hydrogen and oxygen fuels". Proceedings of the National Academy of Sciences. 116 (14): 6624–6629. Bibcode:2019PNAS..116.6624K. doi:10.1073/pnas.1900556116. PMC 6452679. PMID 30886092.
  20. ^ Dresp, Sören; Dionigi, Fabio; Klingenhof, Malte; Strasser, Peter (12 April 2019). "Direct Electrolytic Splitting of Seawater: Opportunities and Challenges". ACS Energy Letters. 4 (4): 933–942. doi:10.1021/acsenergylett.9b00220. S2CID 189716726.
  21. ^ a b c Xu, Shao-Wen; Li, Jianyi; Zhang, Nan; Shen, Wei; Zheng, Yao; Xi, Pinxian (2023). "Recent advances in direct seawater splitting for producing hydrogen". Chemical Communications. 59 (65): 9792–9802. doi:10.1039/d3cc02074f. PMID 37527284. S2CID 260225254.
  22. ^ Khan, M. A.; Al-Attas, Tareq; Roy, Soumyabrata; Rahman, Muhammad M.; Ghaffour, Noreddine; Thangadurai, Venkataraman; Larter, Stephen; Hu, Jinguang; Ajayan, Pulickel M.; Kibria, Md Golam (2021). "Seawater electrolysis for hydrogen production: a solution looking for a problem?". Energy & Environmental Science. 14 (9): 4831–4839. doi:10.1039/d1ee00870f. hdl:10754/670257.
  23. ^ a b Hausmann, J. Niklas; Schlögl, Robert; Menezes, Prashanth W.; Driess, Matthias (2021). "Is direct seawater splitting economically meaningful?". Energy & Environmental Science. 14 (7): 3679–3685. doi:10.1039/D0EE03659E.
  24. ^ Farràs, Pau; Strasser, Peter; Cowan, Alexander J. (August 2021). "Water electrolysis: Direct from the sea or not to be?". Joule. 5 (8): 1921–1923. doi:10.1016/j.joule.2021.07.014.
  25. ^ Maril, Marisol; Delplancke, Jean-Luc; Cisternas, Nataly; Tobosque, Pablo; Maril, Yasmín; Carrasco, Claudia (January 2022). "Critical aspects in the development of anodes for use in seawater electrolysis". International Journal of Hydrogen Energy. 47 (6): 3532–3549. doi:10.1016/j.ijhydene.2021.11.002. S2CID 244561736.
  26. ^ Ping, Qingyun; Cohen, Barak; Dosoretz, Carlos; He, Zhen (September 2013). "Long-term investigation of fouling of cation and anion exchange membranes in microbial desalination cells". Desalination. 325: 48–55. Bibcode:2013Desal.325...48P. doi:10.1016/j.desal.2013.06.025.
  27. ^ Veroneau, Samuel S.; Nocera, Daniel G. (2 March 2021). "Continuous electrochemical water splitting from natural water sources via forward osmosis". Proceedings of the National Academy of Sciences. 118 (9): e2024855118. Bibcode:2021PNAS..11824855V. doi:10.1073/pnas.2024855118. PMC 7936378. PMID 33619109.
  28. ^ Veroneau, Samuel S.; Hartnett, Alaina C.; Thorarinsdottir, Agnes E.; Nocera, Daniel G. (28 February 2022). "Direct Seawater Splitting by Forward Osmosis Coupled to Water Electrolysis". ACS Applied Energy Materials. 5 (2): 1403–1408. doi:10.1021/acsaem.1c03998. S2CID 246661386.
  29. ^ Shi, Le; Rossi, Ruggero; Son, Moon; Hall, Derek M.; Hickner, Michael A.; Gorski, Christopher A.; Logan, Bruce E. (2020). "Using reverse osmosis membranes to control ion transport during water electrolysis". Energy & Environmental Science. 13 (9): 3138–3148. doi:10.1039/d0ee02173c. S2CID 224980142.
  30. ^ Xie, Heping; Zhao, Zhiyu; Liu, Tao; Wu, Yifan; Lan, Cheng; Jiang, Wenchuan; Zhu, Liangyu; Wang, Yunpeng; Yang, Dongsheng; Shao, Zongping (22 December 2022). "A membrane-based seawater electrolyser for hydrogen generation". Nature. 612 (7941): 673–678. Bibcode:2022Natur.612..673X. doi:10.1038/s41586-022-05379-5. PMID 36450987. S2CID 254123372.
  31. ^ a b Blain, Loz (16 March 2022). "Record-breaking hydrogen electrolyzer claims 95% efficiency". New Atlas. Archived from the original on 25 December 2022. Retrieved 25 December 2022.
  32. ^ "2001-High pressure electrolysis – The key technology for efficient H.2" (PDF). Retrieved 25 February 2024.
  33. ^ Ghosh, P.C; Emonts, B; Janßen, H; Mergel, J; Stolten, D (2003). "Ten years of operational experience with a hydrogen-based renewable energy supply system" (PDF). Solar Energy. 75 (6): 469–478. Bibcode:2003SoEn...75..469G. doi:10.1016/j.solener.2003.09.006. Archived from the original (PDF) on 27 March 2009.
  34. ^ J. E. O'Brien; C. M. Stoots; J. S. Herring; M. G. McKellar; E. A. Harvego; M. S. Sohal; K. G. Condie (2010). High Temperature Electrolysis for Hydrogen Production from Nuclear Energy ? TechnologySummary (PDF) (Report). doi:10.2172/978368.
  35. ^ Kalamaras, Christos M.; Efstathiou, Angelos M. (6 June 2013). "Hydrogen Production Technologies: Current State and Future Developments". Conference Papers in Science. 2013: e690627. doi:10.1155/2013/690627.
  36. ^ "High temperature electrolysis using SOEC". Hi2h2. Archived from the original on 3 March 2016. Retrieved 5 May 2016.
  37. ^ "WELTEMPWater electrolysis at elevated temperatures". Weltemp.eu. 31 December 2010. Archived from the original on 3 March 2016. Retrieved 5 May 2016.
  38. ^ Chatenet, Marian; Pollet, Bruno G.; Dekel, Dario R.; Dionigi, Fabio; Deseure, Jonathan; Millet, Pierre; Braatz, Richard D.; Bazant, Martin Z.; Eikerling, Michael; Staffell, Iain; Balcombe, Paul; Shao-Horn, Yang; Schäfer, Helmut (2022). "Water electrolysis: from textbook knowledge to the latest scientific strategies and industrial developments". Chemical Society Reviews. 51 (11): 4583–4762. doi:10.1039/d0cs01079k. PMC 9332215. PMID 35575644.
  39. ^ Carmo, M; Fritz D; Mergel J; Stolten D (2013). "A comprehensive review on PEM water electrolysis". Journal of Hydrogen Energy. 38 (12): 4901. Bibcode:2013IJHE...38.4901C. doi:10.1016/j.ijhydene.2013.01.151.
  40. ^ "Alkaline Water Electrolysis" (PDF). Energy Carriers and Conversion Systems. Retrieved 19 October 2014.
  41. ^ Schalenbach, M; Tjarks G; Carmo M; Lueke W; Mueller M; Stolten D (2016). "Acidic or Alkaline? Towards a New Perspective on the Efficiency of Water Electrolysis". Journal of the Electrochemical Society. 163 (11): F3197. doi:10.1149/2.0271611jes. S2CID 35846371.
  42. ^ a b Jordan, Thomas (2022), "Hydrogen technologies", Hydrogen Safety for Energy Applications, Elsevier, pp. 25–115, doi:10.1016/b978-0-12-820492-4.00005-1, ISBN 978-0-12-820492-4, retrieved 27 April 2024
  43. ^ Alternative electrochemical systems for ozonation of water. NASA Tech Briefs (Technical report). NASA. 20 March 2007. MSC-23045. Retrieved 17 January 2015.
  44. ^ Nakhiah Goulbourne. "Research Topics for Materials and Processes for PEM Fuel Cells REU for 2008". Virginia Tech. Archived from the original on 27 February 2009. Retrieved 18 July 2008.
  45. ^ Zhiwei Yang; et al. (2004). "Novel inorganic/organic hybrid electrolyte membranes" (PDF). Prepr. Pap.-Am. Chem. Soc., Div. Fuel Chem. 49 (2): 599. Archived from the original (PDF) on 28 April 2017. Retrieved 19 October 2021.
  46. ^ US patent 5266421, Townsend, Carl W. & Naselow, Arthur B., "Enhanced membrane-electrode interface", issued 11 November 2008, assigned to Hughes Aircraft 
  47. ^ Gabriel Gache (17 December 2007). "New Proton Exchange Membrane Developed – Nafion promises inexpensive fuel-cells". Softpedia. Archived from the original on 23 April 2008. Retrieved 18 July 2008.
  48. ^ Varcoe, John R.; Atanassov, Plamen; Dekel, Dario R.; Herring, Andrew M.; Hickner, Michael A.; Kohl, Paul. A.; Kucernak, Anthony R.; Mustain, William E.; Nijmeijer, Kitty; Scott, Keith; Xu, Tongwen; Zhuang, Lin (2014). "Anion-exchange membranes in electrochemical energy systems". Energy Environ. Sci. 7 (10): 3135–3191. doi:10.1039/C4EE01303D. hdl:10044/1/24509.
  49. ^ Dekel, Dario R. (January 2018). "Review of cell performance in anion exchange membrane fuel cells". Journal of Power Sources. 375: 158–169. Bibcode:2018JPS...375..158D. doi:10.1016/j.jpowsour.2017.07.117.
  50. ^ "Developing the world's most efficient electrolyser". Supercritical. Archived from the original on 6 November 2021. Retrieved 6 November 2021.
  51. ^ "A low-cost water splitter that runs on an ordinary AAA battery". KurzweilAI. 22 August 2014. Archived from the original on 16 April 2015. Retrieved 11 April 2015.
  52. ^ a b Gong, Ming; Zhou, Wu; Tsai, Mon-Che; Zhou, Jigang; Guan, Mingyun; Lin, Meng-Chang; Zhang, Bo; Hu, Yongfeng; Wang, Di-Yan; Yang, Jiang; Pennycook, Stephen J.; Hwang, Bing-Joe; Dai, Hongjie (2014). "Nanoscale nickel oxide/nickel heterostructures for active hydrogen evolution electrocatalysis". Nature Communications. 5: 4695. Bibcode:2014NatCo...5.4695G. doi:10.1038/ncomms5695. PMID 25146255.
  53. ^ Mulder, F. M.; et al. (2017). "Efficient electricity storage with the battolyser, an integrated Ni-Fe-battery and electrolyser". Energy and Environmental Science. 10 (3): 756–764. doi:10.1039/C6EE02923J. S2CID 99216185. Retrieved 6 June 2020.
  54. ^ Blain, Loz (14 February 2023). ""Exceptional" new catalyst cheaply splits hydrogen from seawater". New Atlas. Archived from the original on 14 February 2023. Retrieved 14 February 2023.
  55. ^ Hodges, Aaron; Hoang, Anh Linh; Tsekouras, George; Wagner, Klaudia; Lee, Chong-Yong; Swiegers, Gerhard F.; Wallace, Gordon G. (15 March 2022). "A high-performance capillary-fed electrolysis cell promises more cost-competitive renewable hydrogen". Nature Communications. 13 (1): 1304. Bibcode:2022NatCo..13.1304H. doi:10.1038/s41467-022-28953-x. PMC 8924184. PMID 35292657.
  56. ^ "Making Space Safer with Electrolysis". ASME. Archived from the original on 15 May 2012. Retrieved 26 May 2012.
  57. ^ "Breathing Easy on the Space Station". NASA Science. Archived from the original on 19 May 2012. Retrieved 26 May 2012.
  58. ^ "It is harder for new electric grids to balance supply and demand". The Economist. ISSN 0013-0613. Archived from the original on 30 April 2023. Retrieved 30 April 2023.
  59. ^ Luca Bertuccioli; et al. (7 February 2014). "Development of water electrolysis in the European Union" (PDF). Client Fuel Cells and Hydrogen Joint Undertaking. Archived from the original (PDF) on 31 March 2015. Retrieved 3 December 2014. (page 10) Archived 10 March 2016 at the Wayback Machine.
  60. ^ Stensvold, Tore (26 January 2016). «Coca-Cola-oppskrift» kan gjøre hydrogen til nytt norsk industrieventyr Archived 5 March 2016 at the Wayback Machine. Teknisk Ukeblad, .
  61. ^ Collins, Leigh (28 April 2022). "'Cheaper green hydrogen' | US start-up's novel low-cost electrolyser promises 30% more bang per buck". Recharge. Archived from the original on 1 May 2022. Retrieved 1 May 2022.
  62. ^ "Helmeth". High Temperature Electrolysis Cell. Archived from the original on 24 May 2022. Retrieved 20 June 2022.
  63. ^ "Lessons learned from SOFC/SOEC Development" (PDF). DOE. Archived (PDF) from the original on 20 June 2022. Retrieved 20 June 2022.
  64. ^ Stolten, Detlef (4 January 2016). Hydrogen Science and Engineering: Materials, Processes, Systems and Technology. John Wiley & Sons. p. 898. ISBN 9783527674299. Archived from the original on 22 April 2018. Retrieved 22 April 2018.
  65. ^ Bernholz, Jan (13 September 2018). "RWE's former, current and possible future energy storage applications" (PDF). RWE. p. 10. Archived from the original on 23 May 2019. Retrieved 23 May 2019. Total Efficiency: 70%, or 86% (usage of waste heat)
  66. ^ "ITM – Hydrogen Refuelling Infrastructure – February 2017" (PDF). level-network.com. p. 12. Archived (PDF) from the original on 17 April 2018. Retrieved 17 April 2018.
  67. ^ "Cost reduction and performance increase of PEM electrolysers" (PDF). Europa (web portal). p. 9. Archived (PDF) from the original on 17 April 2018. Retrieved 17 April 2018.
  68. ^ a b Bjørnar Kruse; Sondre Grinna; Cato Buch (13 February 2002). "Hydrogen—Status and Possibilities" (PDF). The Bellona Foundation. p. 20. Archived from the original (PDF) on 16 September 2013. Efficiency factors for PEM electrolysers up to 94% are predicted, but this is only theoretical at this time.
  69. ^ Blain, Loz (13 May 2024). "World's highest-efficiency hydrogen system scales up for mass production". New Atlas. Retrieved 16 May 2024.
  70. ^ Hodges, Aaron; Hoang, Anh Linh; Tsekouras, George; Wagner, Klaudia; Lee, Chong-Yong; Swiegers, Gerhard F.; Wallace, Gordon G. (15 March 2022). "A high-performance capillary-fed electrolysis cell promises more cost-competitive renewable hydrogen". Nature Communications. 13 (1): 1304. Bibcode:2022NatCo..13.1304H. doi:10.1038/s41467-022-28953-x. PMC 8924184. PMID 35292657.
  71. ^ a b Martin, Polly (27 July 2023). "ANALYSIS | How much does a kilogram of green hydrogen actually cost? Well, it's complicated". rechargenews.com. Power price is the biggest factor in determining how much green H2 costs, accounting for 60-75% of the final cost of hydrogen production.
  72. ^ a b Collins, Leigh (24 April 2024). "ANALYSIS | Clean hydrogen 'remains too expensive and uncompetitive' – how can costs be reduced?". rechargenews.com.
  73. ^ Werner Zittel; Reinhold Wurster (8 July 1996). "Chapter 3: Production of Hydrogen. Part 4: Production from electricity by means of electrolysis". HyWeb: Knowledge – Hydrogen in the Energy Sector. Ludwig-Bölkow-Systemtechnik GmbH. Archived from the original on 7 February 2007. Retrieved 14 January 2006.
  74. ^ "high-rate and high efficiency 3D water electrolysis". Grid-shift.com. Archived from the original on 22 March 2012. Retrieved 13 December 2011.
  75. ^ "DOE Technical Targets for Hydrogen Production from Electrolysis". energy.gov. US Department of Energy. Archived from the original on 23 April 2018. Retrieved 22 April 2018.
  76. ^ Deign, Jason. "Xcel Attracts 'Unprecedented' Low Prices for Solar and Wind Paired With Storage". greentechmedia.com. Wood MacKenzie. Archived from the original on 4 February 2018. Retrieved 22 April 2018.
  77. ^ "Wide Spread Adaption of Competitive Hydrogen Solution" (PDF). nelhydrogen.com. Nel ASA. Archived (PDF) from the original on 22 April 2018. Retrieved 22 April 2018.
  78. ^ Philibert, Cédric. "Commentary: Producing industrial hydrogen from renewable energy". iea.org. International Energy Agency. Archived from the original on 22 April 2018. Retrieved 22 April 2018.
  79. ^ Collins, Leigh (7 March 2022). "Ukraine war | Green hydrogen 'now cheaper than grey in Europe, Middle East and China': BNEF". Recharge. Archived from the original on 5 April 2022.
  80. ^ Collins, Leigh (1 February 2022). "Record breaker | World's largest green hydrogen project, with 150MW electrolyser, brought on line in China". Recharge. Archived from the original on 6 February 2022. Retrieved 6 February 2022.
  81. ^ Heyward, Hack (19 April 2022). "Beijing hydrogen body admits that Chinese electrolysers cannot compete with Western machines — yet". Recharge. Archived from the original on 20 April 2022.
  82. ^ Collins, Leigh (12 April 2022). "8,000% growth | 'More than 100GW of hydrogen electrolysers to be produced annually by 2031'". Recharge. Archived from the original on 12 April 2022.
  83. ^ 1.23 V is the standard potential; in non-standard conditions it may be different, in particular, it decreases with temperature.
  84. ^ Kibsgaard, Jakob; Jaramillo, Thomas F.; Besenbacher, Flemming (2014). "Building an appropriate active-site motif into a hydrogen-evolution catalyst with thiomolybdate [Mo3S13]2− clusters". Nature Chemistry. 6 (3): 248–253. Bibcode:2014NatCh...6..248K. doi:10.1038/nchem.1853. PMID 24557141. Archived from the original on 30 July 2020. Retrieved 1 July 2019.
  85. ^ Fei, Huilong; Ye, Ruquan; Ye, Gonglan; Gong, Yongji; Peng, Zhiwei; Fan, Xiujun; Samuel, Errol L. G.; Ajayan, Pulickel M.; Tour, James M. (2014). "Boron- and Nitrogen-Doped Graphene Quantum Dots/Graphene Hybrid Nanoplatelets as Efficient Electrocatalysts for Oxygen Reduction". ACS Nano. 8 (10): 10837–43. doi:10.1021/nn504637y. PMID 25251218.
  86. ^ Luo, J.; Im, J.-H.; Mayer, M. T.; Schreier, M.; Nazeeruddin, M. K.; Park, N.-G.; Tilley, S. D.; Fan, H. J.; Gratzel, M. (2014). "Water photolysis at 12.3% efficiency via perovskite photovoltaics and Earth-abundant catalysts". Science. 345 (6204): 1593–1596. Bibcode:2014Sci...345.1593L. doi:10.1126/science.1258307. PMID 25258076. S2CID 24613846.
  87. ^ Shwartz, Mark (22 August 2014). "Stanford scientists develop water splitter that runs on ordinary AAA battery". News.stanford.edu. Archived from the original on 16 April 2016. Retrieved 5 May 2016.
  88. ^ "Scientists develop a water splitter that runs on an ordinary AAA battery". Technology.org. 25 August 2014. Archived from the original on 2 April 2016. Retrieved 5 May 2016.
  89. ^ Kondori, Alireza (2 May 2019). "Identifying Catalytic Active Sites of Trimolybdenum Phosphide (Mo3P) for Electrochemical Hydrogen Evolution". Advanced Energy Materials. 9 (22). AdvancedEnergyMaterials: 1900516. Bibcode:2019AdEnM...900516K. doi:10.1002/aenm.201900516. OSTI 1531000.
  90. ^ Shi, Yanmei (25 January 2016). "Recent advances in transition metal phosphide nanomaterials: synthesis and applications in hydrogen evolution reaction". Chemical Society Reviews. 45 (6). ChemicalSocietyReviews: 1529–1541. doi:10.1039/C5CS00434A. PMID 26806563.
  91. ^ Hyman D. Gesser (2002). Applied Chemistry. Springer. pp. 16–. ISBN 978-0-306-46700-4. Retrieved 18 December 2011.
  92. ^ Badwal, Sukhvinder P.S.; Giddey, Sarbjit; Munnings, Christopher (September 2013). "Hydrogen production via solid electrolytic routes". Wiley Interdisciplinary Reviews: Energy and Environment. 2 (5): 473–487. Bibcode:2013WIREE...2..473B. doi:10.1002/wene.50. S2CID 135539661.
  93. ^ Badwal, Sukhvinder P. S.; Giddey, Sarbjit S.; Munnings, Christopher; Bhatt, Anand I.; Hollenkamp, Anthony F. (24 September 2014). "Emerging electrochemical energy conversion and storage technologies (open access)". Frontiers in Chemistry. 2: 79. Bibcode:2014FrCh....2...79B. doi:10.3389/fchem.2014.00079. PMC 4174133. PMID 25309898.
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