Electrolysis of Water
Electrolysis of water is its decomposition to give hydrogen and oxygen gases due to the passage of an electric current.
2H2O + electrical energy O2 + 2H2
Electric effects on water
Magnetic effects on water
Electromagnetic effects on water
Water redox processes
What is less well understood?
Driving cars using water
'I propose to distinguish these bodies by calling those anions which go to the anode ....and those passing to the cathode, cations '
Creating an electric potential through water causes positive ions, including the inherent hydrogen ions H3O+, to move towards the negative electrode (cathode) and negative ions, including the inherent hydroxide ions OH-, to move towards the positive electrode (anode). At sufficient potential difference, this may may cause electrolysis with oxygen gas being produced at the anode and hydrogen gas produced at the cathode (see  for current reviews). The electrolysis of water usually involves dilute, or moderately concentrated, salt solutions in order to reduce the power loss driving the current through the solution, but the presence of salt is not a requirement for electrolysis.
||6H2O(l) O2(g) + 4H3O+(aq) + 4e-(to anode) b
||E° = +1.229 V, pH 0 d
||E°' = +0.815 V
||4e-(from cathode) + 4H2O(l) 2H2(g) + 4OH-(aq)
||E° = -0.828 V, pH 14
||E°' = -0.414 V
where (l), (g) and (aq) show the states of the material as being liquid, gas or aqueous solution and with the electrical circuit passing the electrons back from the anode to the cathode. The reactions are heterogeneous, taking place at the boundary between the electrode and the electrolyte with the aqueous boundary layer subject to concentration and electrical potential gradients with the presence of the generated gaseous nanobubbles and microbubbles.
Generally the water next to the electrodes c will change pH due to the ions produced or consumed. If the electrode compartments are separated by a suitable porous membrane then the concentration of H3O+ in the anolyte and OH- in the catholyte (and hence the increase in the respective conductivities) are both expected to increase more than if there is free mixing between the electrodes, when most of these ions will neutralize each other. Small but expected differences in the solutions’ pHs next to the anode (anolyte) and cathode (catholyte) cause only a slight change to the overall potential difference required (1.229 V). Increasing the acid content next to the anode due to the H3O+ produced will increase its electrode potential (for example: pH 4 E = +0.992 V) and increasing the alkaline content next to the cathode due to the OH- produced will make its electrode potential more negative (for example: pH 10 E = -0.592 V). If the anode reaction is forced to run at pH 14 and the cathode reaction is run at pH 0 then the electrode potentials are +0.401 V and 0 V respectively (see above right). d
(a) Anode pH 0 2 H2O O2 + 4H+ + 4e- E°= +1.229 V
(b) Anode pH 14 4 OH-
O2 +H2O + 4e- E°= +0.401 V
(c) Cathode pH 0 4 H+ + 4e- 2H2 E°= 0.0 V
(d) Cathode pH 14 4 H2O + 4e- 2H2 +4OH- E°= -0.828 V
This does not mean that because the electrolysis can be achieved with a (minimum) voltage of +0.403 V (c and b, above right) , it breaks the thermodynamic requirement of 1.229 V as there is a further input of energy required in keeping the electrode compartments at the required pHs and solute concentration.
The current flowing indicates the rate of electrolysis. The amount of product formed can be calculated directly from the duration and current flowing, as 96,485 coulombs (i.e. one faraday) delivers one mole of electrons; with one faraday ideally producing 0.5 moles of H2 plus 0.25 moles of O2. Thus, one amp flowing for one second (one coulomb) produces 5.18 µmol H2 (10.455 µg, 0.1177 mL at STP) and 2.59 µmol O2 (82.888 µg, 0.0588 mL at STD; 4.9 kW h/m3 H2 at 60% efficiency), if there are no side reactions at the electrodes;
Number of moles = Coulombs/(unsigned numeric charge on the ion ˣ Faraday)
Number of moles = (Current in amperes ˣ time in seconds)/(unsigned numeric charge on the ion ˣ Faraday)
The gases produced at the electrodes may dissolve, with their equilibrium solubility proportional to their partial pressure as gases in the atmosphere above the electrolytic surface. Oxygen gas is poorly soluble (~44 mg kg-1, ~1.4 mM at 0.1 MPa and 20 °C, but only ~0.29 mM against its normal atmospheric partial pressure). Hydrogen gas is less soluble (~1.6 mg kg-1, ~0.80 mM at 0.1 MPa and 20 °C but only ~0.44 nM against its very low normal atmospheric partial pressure). It may take considerable time for the solubilities to drop from their initially-super-saturated state to their equilibrium values after the electrolysis has ended.
Although theoretically as described above, the current passing should determine the amounts of hydrogen and oxygen formed, several factors ensure that somewhat lower amounts of gas are actually found. Some electrons (and product) are used up in side reactions, some of the products are catalytically reconverted to water at the electrodes particularly if there is no membrane dividing the electrolysis compartments, some hydrogen may absorb into the cathode (particularly if palladium is used) and some oxygen oxidizes the anode. Finally some gas remains held up in the nanobubbles for a considerable time and some gas may escape measurement.
The above description hides much important science and grossly over-simplifies the system. The actual potential required at any position within the electrolytic cell is determined by the localized concentration of the reactants and products including the local pH of the solution, instantaneous gas partial pressure and effective electrode surface area loss due to attached gas bubbles. In addition, a greater potential difference (called overpotential) is required at both electrodes to overcome activation energy barriers and deliver a significant reaction rate. Typically at good electrodes, such as those made of platinum, that may total an addition of about half a volt to the potential difference between the electrodes. In addition a further potential difference is required to drive the current through the electrical resistance of electrolytic cell and circuit; for a (typical) one ohm cell circuit resistance, a each amp current flow would require a further one volt and waste one watt of power. This power (and consequent energy) loss (~20%, ) causes the electrolyte to warm up during electrolysis.
The minimum necessary cell voltage to start water electrolysis is the potential 1.229 V.
The potential necessary to start water electrolysis without withdrawing heat from the surroundings is
-ΔH°'/nF = 1.481 V.
This results in at least a 21% unavoidable loss of efficiency. Normally further heat is generated, and efficiency lost, from the overpotentials applied.
The efficiency of electrolysis  increases with the temperature as the hydrogen bonding reduces. If the pressure over the electrolysis is increased then more current passes for the same applied voltage. However the output of gas per coulomb and the heating effect are both decreased. This is due to the increased solubility of the gases and smaller bubbles both reducing the cell resistance and increasing recombination reactions. Although reducing the distance between electrodes reduces the resistance of the electrolysis medium, the process may suffer if the closeness allows a build up of gas between these electrodes . Low to higher pulsed potential increases the reaction (current) and accelerates both the movement of bubbles from the electrode surface and the mass transfer rate in the electrolyte, which lowers the electrochemical polarization in the diffusion layer and further increases hydrogen production efficiency . The rate of change of the current density (and hence efficiency) can be increased using a magnetic field .
Pure water conducts an electric current very poorly and, for this reason, is difficult (slow) to electrolyze. However, usually some salts will be added or present in tap and ground waters which will be sufficient to allow electrolysis to proceed at a significant rate. However such salts, and particularly chloride ions, may then undergo redox reactions at an electrode. These side reactions both reduce the efficiency of the electrolysis reactions (above) and produce new solutes. Other electrolytic reactions may occur at the electrodes so producing further solutes and gases. In addition, these solutes may react together to produce other materials. Together the side reactions are complex and this complexity increases somewhat when the voltage applied to the cell is greater than that required by the above reactions and processes. The likely reactions within the electrode compartments are described below. Some of these may only occur to a very small extent and other reactions may also be occurring that are not included. Standard electrode potentials are shown.
Right is given a representation of the compartments in the electrolytic cell with some of their constituent molecules, ions and radicals. Other materials may be present and some of the materials given may be at very low concentrations and/or have short half-lives.
Important amongst the side products is ozone (O3, see left).
The relative amount of O3 produced (relative to molecular oxygen) depends on the overpotential, pH, radicals present and anode material. O2 evolution is greater than that for O3 due to the lower potential required. At low overpotentials, very little O3 may be produced but at high current densities and overpotential, up to a sixth (or more) of the oxidized molecules may be O3. As O3 is more much more soluble than O2, there may twice the dissolved O3 than O2 but the bubble gas will contain about 20 times the O2 than O3 . Tin oxide anodes have proved useful for the production of O3, particularly if doped with Sb and Ni, as they bind both oxygen molecules and hydroxyl radicals to facilitate the O3 production . Decomposition of ozone gives rise to several strong oxidants including hydroxyl radicals (·OH); an extremely strong oxidizer capable of killing viruses, amoebae, algae and dangerous bacteria, such as MRSA and Legionella.
Although charged ions are attracted into the compartments by virtue of the applied potential, oppositely charged ions are created in both compartments due to the electrolytic reactions. Thus for example, Na+ ions enter the catholyte from the anode compartment but excess OH- is produced at the same time at the cathode. The concentration of the OH- ions will be generally expected to be greater than the increase in cations in the catholyte and the concentration of the H3O+ ions will be generally expected to be greater than any increase in anions in the anolyte. Often a conductive but semi-permeable membrane (nafion, a strongly hydrated sulfonated tetrafluoroethylene based copolymer , for example) is used to separate the two compartments and reduce the movement of the products between the electrode compartments; a process that improves the yield by reducing back and side reactions . Due to the easier electrolysis of water containing 1H rather than 2H (D) or 3H (T), electrolysis can be used for producing water with reduced or enriched isotopic composition.
What is less well understood?
Although much time has been spent on investigating and modeling the electrolytic system , it is still not entirely clear how water is arranged at the surface of the electrodes. Alignment of the water dipoles with the field is expected, together with the consequential breakage of a proportion of the water molecules’ hydrogen bonds. When the electrode processes occur it appears that singly-linked hydrogen atoms and singly-linked oxygen atoms are bound to the platinum atoms at the cathode and anode respectively. These bound atoms are able to diffuse around in two dimensions on the surface of their respective electrodes until they take part in further reaction. Other atoms and polyatomic groups may also bind similarly to the electrode surfaces and subsequently undergo reactions. Molecules such as O2 and H2 produced at the surfaces may enter nanoscopic cavities in the liquid water (nanobubbles) as gases or become solvated by the water.
Gas-containing cavities in liquid solution (often called bubbles) grow or shrink by diffusion according to whether the solution is over-saturated or under-saturated with the dissolved gas. Given suitable electrodes, the size of the cathodic hydrogen bubbles depends on the overvoltage with nanobubbles being formed at low overvoltages and larger bubbles being formed at higher overvoltages . Larger micron-plus sized bubbles have sufficient buoyancy to rise through the solution and release contained gas at the surface before all the gas dissolves. With smaller bubbles a pressure is exerted by the surface tension is in inverse proportion to their diameter and they may be expected to collapse. However, as the nanobubble gas/liquid interface is charged, an opposing force to the surface tension is introduced, so slowing or preventing their dissipation. Electrolytic solutions have been proven to contain very large numbers of gaseous nanobubbles . The ‘natural’ state of such interfaces appears to be negative . Other ions with low surface charge density (such as Cl-, ClO-, HO2- and O2·-) will also favor the gas/liquid interfaces [928a] as probably do hydrated electrons [1841, 1874]. Aqueous radicals also prefer to reside at such interfaces . From this known information it seems clear that the nanobubbles present in the catholyte will be negatively charged but those in the anolyte  will probably possess little charge (with the produced excess positive H3O+ ions canceling out the natural negative charge). Therefore, catholyte nanobubbles are not likely to lose their charge on mixing with the anolyte stream and are otherwise known to be stable for many minutes . Additionally gas molecules may become charged within the nanobubbles (such as the superoxide radical ion, O2·-), due to decay of ozone present and the excess potential on the cathode, increasing the overall charge of the nanobubbles and, probably, the stability of that charge. Raised temperature at the electrode surface, due to the excess power loss over that required for the electrolysis, may also increase nanobubble formation by reducing local gas solubility. Clearly increasing the pressure on solutions containing nanobubbles will also slow down their dissipation if this pressure has increased the dissolved gas content.
Driving cars using water?
The hydrogen produced by electrolysis may be used as a fuel in a fuel cell (see right) but the efficiency of the overall process (synthesis of H2 from H2O followed by oxidation of H2 to H2O) is always well below 100%. Thus the hydrogen produced can never be used to drive the electrolysis that produces it . This fact is governed by the unbreakable laws of thermodynamics but often seems to be ignored by people proposing cars that run on 'water' or 'Brown's gas' (a mixture of H2 and O2 produced by electrolysis ). Generating hydrogen by electrolysis is only (optimally) about 60% efficient and the use of this hydrogen in a car is (optimally) also only about 60% efficient, so two thirds of the energy required is wasted. The only time excess energy may apparently be produced (on a laboratory scale) is when the electrodes themselves react; a important factor that produces artifacts when using some stainless steel electrodes but often ignored.
An alternative source of energy in water electrolysis has been proposed . It involves 'cold fusion' (low energy nuclear reactions), a highly controversial theory developed from experiments involving electrolysis of heavy water using palladium (Pd) cathodes and reportedly producing greater heat than could be conventionally explained. This idea has so far received only limited acceptance with the main criticism being a lack of a suitable theoretical basis. History will decide. [Back to Top ]
a The approximate resistivities of pure water, tap water and sea water are 18 MΩ cm, 5 kΩ cm and 20 Ω cm respectively. Thus the rate of the electrolysis is speeded up by factors of about 1000 ˣ or 1,000,000 ˣ using tapwater or seawater respectively rather than pure water. [Back]
b Traditionally such equations are written with the electrons on the left-hand-side and (however written) the redox potential refers to so directed equations. Here it is written reversed in order to show how the cell reaction is balanced, as this is the way the reaction occurs.
| O2(g) + 4H3O+(aq) + 4e- 6H2O(l)
||E° = +1.229 V, pH 0
c The electrodes should preferably be made from material with high conductivity, resistance to corrosion and erosion during the electrolysis and able to catalyze the electrode reactions. Also for industrial use, they should be relatively inexpensive. Platinum is an excellent but expensive electrode material. Industrial cathodes may be made from steel or nickel and those used as anodes are metals such as titanium coated with the oxides and mixed oxides of metals such as nickel and cobalt. Water next to the surface will organize dependent on the surface material . [Back]
d At the anode, E° = +1.229 - 0.059 pH V At the cathode E° = - 0.059 pH V. The value '0.059' is derived from the Nernst constant = Loge(10) ˣ RT/F = 0.059 V (25 °C). [Back]