Inchemistry andmanufacturing,electrolysis is a technique that usesdirect electric current (DC) to drive an otherwise non-spontaneouschemical reaction. Electrolysis is commercially important as a stage in the separation ofelements from naturally occurring sources such asores using anelectrolytic cell. Thevoltage that is needed for electrolysis to occur is called the decomposition potential. The word "lysis" means to separate or break, so in terms, electrolysis would mean "breakdown via electricity."
The word "electrolysis" was introduced byMichael Faraday in 1834,[1] using theGreek wordsἤλεκτρον[ɛ̌ːlektron] "amber", which since the 17th century was associated withelectrical phenomena, andλύσις[lýsis] meaning "dissolution". Nevertheless, electrolysis, as a tool to study chemical reactions and obtain pureelements, precedes the coinage of the term and formal description by Faraday.
In the early nineteenth century,William Nicholson andAnthony Carlisle sought to furtherVolta's experiments. They attached two wires to either side of avoltaic 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 namedMartin 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 1791Luigi 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 toHumphry 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 1817Johan 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 thatWilliam 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 twolaws of electrolysis.[5]
On June 26, 1886,Ferdinand Frederick Henri Moissan finally felt comfortable performing electrolysis onanhydrous 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 theAcadémie 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 1886Charles Martin Hall from America andPaul 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 theHall–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 ofcopper andzinc.[16][17][18]
1821 –Lithium was discovered by the English chemistWilliam Thomas Brande, who obtained it by electrolysis of lithium oxide.
1834 –Michael Faraday published histwo laws of electrolysis, provided a mathematical explanation for them, and introduced terminology such as electrode, electrolyte, anode, cathode, anion, and cation.
1902 – Stanisław Łaszczyński obtainedcopper using electrolysis.[16][17][18]
1930 – Development of the modern chlor-alkali process (electrolysis of brine to produce chlorine and sodium hydroxide), which became an important industrial method.[21]
The main components required to achieve electrolysis are anelectrolyte, electrodes, and an external power source. A partition (e.g. anion-exchange membrane or asalt bridge) is optional to keep the products from diffusing to the vicinity of the opposite electrode.
The electrolyte is achemical substance which containsfree ions and carrieselectric current (e.g. an ion-conductingpolymer, solution, or aionic liquid compound). If the ions are not mobile, as in most solidsalts, 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 theelectrolyte and are connected to the power source which completes theelectrical circuit. Adirect current supplied by the power source drives the reaction causing ions in the electrolyte to be attracted toward the respective oppositely charged electrode.
Electrodes ofmetal,graphite andsemiconductor material are widely used. Choice of suitableelectrode 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.[22] 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.
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 theirequivalent weight. These are known asFaraday's laws of electrolysis.
Each electrode attracts ions that are of the oppositecharge. Positively charged ions (cations) move towards the electron-providing (negative) cathode. Negatively charged ions (anions) move towards the electron-extracting (positive) anode. In this processelectrons are effectively introduced at the cathode as areactant and removed at the anode as aproduct. In chemistry, the loss of electrons is calledoxidation, while electron gain is calledreduction.
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 areelectroplating,electrowinning, andelectrorefining.
When an ion gains or loses electrons without becoming neutral, its electronic charge is altered in the process.
For example, theelectrolysis of brine produces hydrogen and chlorine gases which bubble from the electrolyte and are collected. The initial overall reaction is thus:[23]
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 morebasic due to the production of OH−, less Cl2 emerges from the solution as it begins to react with the hydroxide producinghypochlorite (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 theself-ionization of water and the decomposition of hypochlorite at the cathode, the rate of the latter depends on factors such asdiffusion and the surface area of the cathode in contact with the electrolyte.[24]
Decomposition potential or decomposition voltage refers to the minimum voltage (difference inelectrode potential) betweenanode andcathode of an electrolytic cell that is needed for electrolysis to occur.[25]
The voltage at which electrolysis is thermodynamically preferred is the difference of the electrode potentials as calculated using theNernst equation. Applying additional voltage, referred to asoverpotential, can increase the rate of reaction and is often needed above the thermodynamic value. It is especially necessary for electrolysis reactions involving gases, such asoxygen,hydrogen orchlorine.
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 agas diffusion electrode.
The amount of electrical energy that must be added equals the change inGibbs free energy of the reaction plus the losses in the system. The losses can (in theory) be arbitrarily close to zero, so the maximumthermodynamic efficiency equals theenthalpy 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 ofsteam into hydrogen and oxygen at high temperature, the opposite is true and heat energy is absorbed. This heat is absorbed from the surroundings, and theheating value of the produced hydrogen is higher than the electric input.
Pulsating current results in products different from DC. For example, pulsing increases the ratio ofozone to oxygen produced at the anode in the electrolysis of an aqueous acidic solution such as dilute sulphuric acid.[26] Electrolysis of ethanol with pulsed current evolves an aldehyde instead of primarily an acid.[27]
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.
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 5–6V. Theanode, the electrocatalyst, isnickel-plated.
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.
InElectrochemical 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 fordeburring or for etching metal surfaces like tools or knives with a permanent mark or logo.
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 sulfateSO2− 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 thestandard electrode potential for the reactions at the anode and cathode. The standard electrode potential is directly related to theGibbs free energy, ΔG, for the reactions at each electrode and refers to an electrode with no current flowing. An extract from thetable of standard electrode potentials is shown below.
In terms of electrolysis, this table should be interpreted as follows:
Movingdown the table,E° becomes more positive, and species on theleft are more likely to bereduced: for example, zinc ions are more likely to be reduced to zinc metal than sodium ions are to be reduced to sodium metal.
Movingup the table,E° becomes more negative, and species on theright are more likely to beoxidized: for example, sodium metal is more likely to be oxidized to sodium ions than zinc metal is to be oxidized to zinc ions.
Using theNernst equation theelectrode 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 Mzinc 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.[34]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 termedoverpotentials. 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 ofchloride tochlorine is lower than the overpotential for the oxidation ofwater tooxygen. Thehydroxide ions and dissolved chlorine gas react further to formhypochlorous acid. The aqueous solutions resulting from this process is calledelectrolyzed water and is used as a disinfectant and cleaning agent.
The electrochemical reduction or electrocatalytic conversion ofCO2 can produce value-added chemicals such asmethane,ethylene,ethanol, etc.[35][36][37] The electrolysis of carbon dioxide gives formate or carbon monoxide, but sometimes more elaborate organic compounds such asethylene.[38] This technology is under research as a carbon-neutral route to organic compounds.[39][40]
Theenergy 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);[41] however, higher practical efficiencies are available with the use ofpolymer electrolyte membrane electrolysis and catalytic technology, such as 95% efficiency.[42][43]
TheNational 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.[44]
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 theHaber process, and converting heavy petroleum sources to lighter fractions viahydrocracking. Onsite electrolysis has been utilized to capture hydrogen for hydrogen fuel-cells inhydrogen vehicles.
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.[45][46] The carbon/hydrocarbon assisted water electrolysis (so-called CAWE) process for hydrogen generation would perform this operation in a singleelectrochemical reactor. This system energy balance can be required only around 40% electric input with 60% coming from the chemical energy of carbon or hydrocarbon.[47] 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:[48]
Carbon/Coal slurry (C + 2H2O) → CO2 + 2H2E′ = 0.21 V (reversible voltage) /E′ = 0.46 V (thermo-neutral voltage)
or
Carbon/Coal slurry (C + H2O) → CO + H2E′ = 0.52 V (reversible voltage) /E′ = 0.91 V (thermo-neutral voltage)
Thus, this CAWE approach is that the actual cell overpotential can be significantly reduced to below 1.0 V as compared to 1.5 V for conventional water electrolysis.
A specialized application of electrolysis involves the growth of conductive crystals on one of the electrodes from oxidized or reduced species that are generated in situ. The technique has been used to obtain single crystals of low-dimensional electrical conductors, such ascharge-transfer salts andlinear chain compounds.[49][50]
The current method of producing steel fromiron 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 ofCO2e with 22% of that being direct emissions from the blast furnace.[51] As of 2022, steel production contributes 7–9% of global emissions.[52] 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 moltenoxide salts and using a platinum anode.[53] 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, anodecorrosion is a concern at these temperatures.
Additionally, the low temperature reduction of iron oxide by dissolving it in alkaline water has been reported.[54] 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.
A 2020 study investigated direct electrolysis of seawater, alkaline electrolysis,proton-exchange membrane electrolysis, andsolid oxide electrolysis.[55] 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 theanode,2 Cl−(aq) → Cl2(g) + 2e− and the reaction at thecathode,2 H2O(l) + 2 e− → H2(g) + 2OH−(aq). The inclusion ofmagnesium andcalciumions in the seawater makes the production of alkalihydroxides 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 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 the simple electrolyzer above are there. There are micro-electrolyzer designs that are able to eliminate the separator requirement by designing the internal flow to separate the gases autonomously. See for example US12116679B2 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 reactionsO−2(g) → 1/2 O2(g) + 2 e− at the anode andH2O(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 ofmultiple-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.
^The Supplement (1803 edition) to Encyclopædia Britannica 3rd edition (1797), volume 1, page 225, "Mister Van Marum, by means of his great electrical machine, decomposed the calces of tin, zinc, and antimony, and resolved them into their respective metals and oxygen" and gives as a reference Journal de Physiques, 1785.
^Blum, W.; Vinal, G. W. (1934). "The Definition of Polarization, Overvoltage and Decomposition Potential".Transactions of the Electrochemical Society.66: 359.doi:10.1149/1.3498105.
^Over, Herbert (2012). "Surface Chemistry of Ruthenium Dioxide in Heterogeneous Catalysis and Electrocatalysis: From Fundamental to Applied Research".Chemical Reviews.112 (6):3356–3426.doi:10.1021/cr200247n.PMID22423981.
^Landolt, D.; Ibl, N. (1972). "Anodic Chlorate Formation on Platinized Titanium".Journal of Applied Electrochemistry.2 (3). Chapman and Hall Ltd.:201–210.doi:10.1007/BF02354977.S2CID95515683.
^Lee, Seunghwa; Ju, Hyungkuk; Machunda, Revocatus; Uhm, Sunghyun; Lee, Jae Kwang; Lee, Hye Jin; Lee, Jaeyoung (2015). "Sustainable production of formic acid by electrolytic reduction of gaseous carbon dioxide".J. Mater. Chem. A.3 (6): 3029.doi:10.1039/C4TA03893B.S2CID98110035.
^Whipple, Devin T.; Kenis, Paul J.A. (2010). "Prospects of CO2 Utilization via Direct Heterogeneous Electrochemical Reduction".The Journal of Physical Chemistry Letters.1 (24): 3451.doi:10.1021/jz1012627.S2CID101946630.
^Machunda, Revocatus L.; Ju, Hyungkuk; Lee, Jaeyoung (2011). "Electrocatalytic reduction of CO2 gas at Sn based gas diffusion electrode".Current Applied Physics.11 (4): 986.Bibcode:2011CAP....11..986M.doi:10.1016/j.cap.2011.01.003.
^Hori, Y (2008). "Electrochemical CO2 Reduction on Metal Electrodes". In C.G. Vayeanas, R. White and M.E. Gamboa-Aldeco (ed.).Modern Aspects of Electrochemistry. Vol. 42 (42 ed.). New York: Springer. pp. 141–153.doi:10.1007/978-0-387-49489-0_3.ISBN978-0-387-49488-3..
^Qiao, Jinli; Liu, Yuyu; Hong, Feng; Zhang, Jiujun (2014). "A review of catalysts for the electroreduction of carbon dioxide to produce low-carbon fuels".Chem. Soc. Rev.43 (2):631–675.doi:10.1039/C3CS60323G.PMID24186433.
^Ju, Hyungkuk; Giddey, Sarbjit; Badwal, Sukhvinder P.S.; Mulder, Roger J. (2016). "Electro-catalytic conversion of ethanol in solid electrolyte cells for distributed hydrogen generation".Electrochimica Acta.212:744–757.doi:10.1016/j.electacta.2016.07.062.
^Ju, Hyungkuk; Giddey, Sarbjit; Badwal, Sukhvinder P.S. (2018). "Role of iron species as mediator in a PEM based carbon-water co-electrolysis for cost-effective hydrogen production".International Journal of Hydrogen Energy.43 (19):9144–9152.Bibcode:2018IJHE...43.9144J.doi:10.1016/j.ijhydene.2018.03.195.
^Bechgaard, K.; Carneiro, K.; Rasmussen, F. B.; Olsen, M.; Rindorf, G.; Jacobsen, C.S.; Pedersen, H.J.; Scott, J.C. (1981). "Superconductivity in an organic solid. Synthesis, structure, and conductivity of bis(tetramethyltetraselenafulvalenium) perchlorate, (TMTSF)2ClO4".Journal of the American Chemical Society.103 (9): 2440.Bibcode:1981JAChS.103.2440B.doi:10.1021/ja00399a065.
