Electrolysis of water

Electrolysis of water is using electricity to split water into oxygen and hydrogen 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. Most of the time, hydrogen is made by splitting methane into carbon dioxide and hydrogen via steam reforming. This is a carbon-intensive process that means for every kilogram of "grey" hydrogen produced, approximately 10 kilograms of CO₂ are emitted into the atmosphere.

History

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. 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. 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.

Principles

A DC electrical power source is connected to two electrodes, or two plates that are placed in the water. Hydrogen appears at the cathode, and oxygen at the anode. 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. However, in many cells competing side reactions occur, resulting in additional products and less than ideal faradaic efficiency.
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.
Efficiency is increased through the addition of an electrolyte and electrocatalysts.

Equations

In pure water at the negatively charged cathode, a reduction reaction takes place, with electrons 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:
In the presence of base, the equations are:
Combining either half reaction pair yields the same overall decomposition of water into oxygen and hydrogen:
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

The decomposition of pure water into hydrogen and oxygen at standard temperature and pressure is not favorable in thermodynamic terms.
Thus, the standard potential of the water electrolysis cell is −1.229 V at 25 °C at pH 0. At 25 °C with pH 7, 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°. For two water molecules electrolysed and hence two hydrogen molecules formed, n = 4, and

ΔG° = 474.48 kJ/2 mol = 237.24 kJ/mol
ΔS° = 163 J/K mol
ΔH° = 571.66 kJ/2 mol = 285.83 kJ/mol
ΔH° = 141.86 kJ/g.

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

Electrolyte

Electrolysis in pure water consumes/reduces H⁺ cations at the cathode and consumes/oxidizes hydroxide 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 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 since it has a low autoionization, K_w = 1.0×10⁻¹⁴ at room temperature and thus pure water conducts current poorly, 0.055 μS/cm. 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.
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, 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, as it is difficult to oxidize. The standard potential for oxidation of this ion to the peroxydisulfate ion is +2.010 volts.
Strong acids such as sulfuric acid, and strong bases such as potassium hydroxide, and sodium hydroxide 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.

Pure water

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, 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, 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.

Seawater

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. Multiple methods have been advanced for electrolysing unprocessed seawater. Typical proton exchange membrane 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. 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.
However, indirect seawater electrolysis plants require more space, energy, and more maintenance, and some believe that the water purity achieved through seawater reverse osmosis may not be sufficient, necessitating additional equipment and cost. 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. 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.
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.
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. Some studies have explored the use of low-cost reverse osmosis membranes to replace expensive ion exchange membranes. 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.
An alternative method involves employing a hydrophobic membrane to prevent ions from entering the cell stack. This method combines a hydrophobic porous polytetrafluoroethylene 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.