electrolysis


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Electrolysis is a chemical process in which an electric current forces a redox reaction . It is used, for example, to extract metals or to manufacture substances which would be more expensive or hardly possible to extract using purely chemical processes. Examples of important electrolyses are the production of hydrogen , aluminum , chlorine and caustic soda .

Electrolysis requires a DC voltage source , which supplies the electrical energy and drives the chemical reactions. Part of the electrical energy is converted into chemical energy . Batteries , accumulators or fuel cells serve exactly the opposite purpose, the conversion of chemical energy into electrical energy : they serve as a power source. When you charge an accumulator, electrolysis takes place, which reverses the chemical processes during the discharge. Electrolyses can therefore serve to store energy, for example in the electrolysis of water , the hydrogen andOxygen results, which have been proposed as a fuel in a hydrogen economy. By reversing the electrolysis of water in a fuel cell, around 40% of the electrical energy originally used can be recovered.

The deposition of metals from a solution containing the corresponding metal ions by an externally applied current is also an electrolysis. This can be used to produce metal layers, for example when chrome plating ; this type of electrolysis is the subject of electroplating . The electrolytic dissolution and redeposition of metals is used for cleaning, e.g. B. of copper , and is called electrolytic refining .

In the chemical reactions that take place during electrolysis, electrons are transferred. There are therefore always redox reactions , with the oxidation taking place at the anode ( electrical pole ) and the reduction taking place at the cathode ; Oxidation and reduction processes are spatially at least partially separated from one another.

history

The electrolysis was discovered in 1800, wherein the by Alessandro Volta invented first useful battery was used, the voltaic pile . The newly discovered electrolysis enabled Humphry Davy to produce several base metals in elemental form for the first time in 1807 and 1808, for example sodium and calcium . Michael Faraday examined electrolysis more closely and discovered its basic laws , namely the dependence of the converted masses on the amount of charge and the molar mass . At his suggestion, the terms electrolysis, electrode , electrolyte , anode , cathode , anion and cation were created. After the invention of powerful electric generators , electrolysis led to a rapid development in science and technology at the end of the 19th century, e.g. B. in the electrolytic production of aluminum , chlorine and alkalis , and in explaining the behavior of electrolytes, which also include acids and bases .

principle

Electrolysis (general)
Example of electrolysis with a zinc iodide solution (any electrode material)

An electrical direct current is passed through two electrodes into a conductive liquid (see electrolyte ). The electrolysis produces reaction products on the electrodes from the substances contained in the electrolyte.

The voltage source causes a shortage of electrons in the electrode connected to the positive pole ( anode ) and an excess of electrons in the other electrode connected to the negative pole ( cathode ). The solution between the cathode and anode contains positively and negatively charged ions as electrolytes. Positively charged ions (cations) or electronically neutral substances take up electrons at the cathode and are thereby reduced. The opposite process takes place at the anode, the release of electrons into the electrode, whereby substances, e.g. B. anions, are oxidized. The amount of electrons transferred at the anode is the same as that transferred at the cathode.

The substances are transported to the electrodes by convective mass transfer (diffusion within the liquid with superimposed flow of the liquid) and, as far as ions are concerned, additionally by migration (migration due to the action of the electric field between the electrodes).

The minimum voltage that must be applied for electrolysis is known as the decomposition voltage (U z or E z ). This voltage or a higher voltage must be applied in order for the electrolysis to take place at all. If this minimum voltage is not reached, the electrolyte or its interfaces with the electrodes, which are also known as the electrochemical double layer , have an insulating effect.

For every substance, for every conversion of ions to di- or polyatomic molecules, the decomposition voltage and the deposition potential can be determined based on the redox potential . The redox potential provides further information, such as the electrolytic decomposition of metal electrodes in acid or the reduction of the decomposition voltage by changing the pH value . The redox potential can be used to calculate that the anodic oxygen formation during the electrolysis of water in a basic solution (decomposition voltage: 0.401 V) takes place at a lower voltage than in an acidic (decomposition voltage: 1.23 V) or neutral (decomposition voltage: 0.815 V) solution At the cathode, on the other hand, hydrogen forms more easily under acidic conditions than under neutral or basic conditions.

If there are several reducible cations in an electrolyte solution, the cations that have a more positive (weaker negative) potential in the redox series (voltage series) are reduced first . When an aqueous saline solution is electrolyzed, hydrogen and not sodium is normally formed at the cathode. Even if there are several types of anions that can be oxidized, those that are as close as possible to zero voltage in the redox series, i.e. have a weaker positive redox potential, come into play.

After the decomposition voltage is exceeded, the current strength increases proportionally with the increase in voltage. According to Faraday, the weight of an electrolytically formed substance is proportional to the electrical charge that has flowed (current strength multiplied by time, see Faraday's laws ). An electrical charge of 96485 C (1 C = 1 A · s) is required for the formation of 1 g of hydrogen (about 11.2 liters, two electrons are required to form a hydrogen molecule) from aqueous solution  . With a current of 1 A, the formation of 11.2 liters of hydrogen takes 26 hours and 48 minutes.

In addition to the redox potential, the overvoltage (the overpotential) is also important. Due to kinetic inhibitions on electrodes, a significantly higher voltage is often required than can be calculated from the calculation of the redox potential. The overvoltage effects can - depending on the material properties of the electrodes - also change the redox series, so that other ions are oxidized or reduced than would have been expected based on the redox potential.

Shortly after the electrolysis has been switched off, an ammeter can be used to detect a current surge in the other direction. In this short phase, the reverse process of electrolysis, the formation of a galvanic cell, begins. In this case, electricity is not used for the implementation, but electricity is generated for a short time; this principle is used in fuel cells .

Sometimes it is advisable to separate the cathode compartment and anode compartment from each other to avoid undesirable chemical reactions and to allow the charge exchange between anode and cathode compartment to take place only through a porous diaphragm - often an ion exchange resin. This is very important in technical electrolysis for the production of caustic soda . Knowledge of molar limit conductivities can also be important for tracking the metabolism and migration speeds of ions .

If you force a separation of individual molecules or bonds through electrolysis, a galvanic element acts at the same time, the voltage of which counteracts the electrolysis. This voltage is also known as the polarization voltage .

Electrodes

There are only a few anode materials that remain inert during electrolysis, i.e. do not go into solution, e.g. B. platinum and carbon . Some metals do not dissolve in spite of their strongly negative redox potential, this property is called "passivity". In acidic solution, according to Nernst's equation, the majority of metals would have to dissolve with the formation of cations and hydrogen. With the exception of copper, silver, gold, platinum and palladium, almost all metal / metal cation pairs have a negative redox potential and would be unsuitable for electrolysis in an acidic environment, as the equilibrium (metal atom and protons) for cation formation and hydrogen is shifted. In the sulfuric acid environment, lead is an inexpensive and popular cathode material; both lead and lead oxide can be used as an anode (also used in car batteries). Lead sulfate is poorly soluble, so the lead electrodes hardly dissolve.

Because of their passivity, iron and nickel can sometimes also be used as anodes in an acidic medium, but these anode materials are also preferably used in a basic medium. An iron anode that has been treated with concentrated nitric acid does not dissolve; no iron (II) or iron (III) ions go into solution through passivation. A very thin and stable iron oxide layer (similar to aluminum) has formed, which prevents the electrode from further dissolving. However, chloride ions or higher temperatures can cancel out passivity.

Compared to other anode materials, iron anodes have only a very low overvoltage in the development of oxygen, which is why they are preferably used in the production of oxygen.

Inhibition phenomena at the anode, which lead to overvoltage during oxygen formation, are observed with carbon and platinum anodes. The overvoltage can be used to generate chlorine instead of oxygen during the electrolysis of aqueous saline solution.

On zinc, lead and especially mercury cathodes, protons show a considerable overvoltage and the formation of hydrogen only takes place at a much higher voltage. The considerable overvoltage of hydrogen at the mercury cathode, to which sodium is bound as an amalgam and thus removed from the equilibrium, is used for the technical production of caustic soda. Due to the considerable overvoltage on this electrode during hydrogen formation, the redox series changes; instead of protons, sodium ions are now reduced on the mercury cathode.

Suitable electrode materials:

metal Used as a cathode Used as an anode Frequent electrolyses
Graphite (burned) + + + + Melt electrolysis (Na, Li, Ca)
Graphite (burned) - + + Aluminum electrolysis
Carbon (smooth) - + Fluorine production
platinum + - Persulfuric acid
iron + + + Water electrolysis
iron + - Melt electrolysis (Na, Li, Ca)
Lead-silver alloy - + Low oxygen overvoltage / fuel cell
lead - + Electrolysis in sulfuric acid solution
lead - + Perchloric acid
aluminum + - Zinc, cadmium electrolysis
Titanium (Ru) - + + Highly resistant to NaCl electrolysis
mercury + - Alkali electrolysis
Tin + copper + - Organic compounds

(++) Well suited, (+) suitable, (-) not suitable

Overload

Overvoltages can occur at the cathode as well as at the anode and thus increase the required voltage compared to the calculations using the Nernst equation . The overvoltages are sometimes considerable in the case of gas formation (e.g. hydrogen and oxygen formation). The applied overvoltage energy is lost as heat, so it does not contribute to the metabolism. The overvoltages vary depending on the type of metal and surface properties of the electrodes. Current strength and temperature also influence the overvoltage. An increasing current strength slightly increases the overvoltage, while an increase in temperature lowers the overvoltage.

The following tables provide a brief overview of the overvoltage in the anodic evolution of oxygen and the cathodic evolution of hydrogen (the tests were carried out at different pH values, for the calculation of pH changes see Nernst equation )

Overvoltage oxygen formation

Conditions: 1 N-aq. KOH, 20 ° C, measurement after 20 min.

Current / area Current / area Current / area
0.01 A / cm 2 0.1 A / cm 2 1 A / cm 2
metal Voltage (V) Voltage (V) Voltage (V)
copper 0.66 0.73 0.77
silver 0.71 0.94 1.06
gold 1.05 1.53 1.63
iron 0.48 0.56 0.63
graphite 0.96 1.12 2.20
nickel 0.75 0.91 1.04
platinum 1.32 1.50 1.55
palladium 1.01 1.12 1.28
lead 0.97 1.02 1.04

Overvoltage hydrogen formation

Conditions: 1 N aq. HCl, 16 ° C.

Current / area Current / area Current / area
0.01 A / cm 2 0.1 A / cm 2 1 A / cm 2
metal Voltage (V) Voltage (V) Voltage (V)
copper 0.75 0.82 0.84
silver 0.66 0.76 -
gold 0.25 0.32 0.42
iron 0.53 0.64 0.77
graphite 0.76 0.99 1.03
nickel 0.42 0.51 0.59
platinum 0.35 0.40 0.40
Platinum-plated platinum 0.03 0.05 0.07
lead 1.24 1.26 1.22
tin 0.98 0.99 0.98
mercury 1.15 1.21 1.24
tungsten 0.35 0.47 0.54

With other electrolytic reductions (without gas formation) the diffusion overvoltage can also become important. If after a few minutes the concentration of the electrolytically converted substance in front of the electrode drops, more voltage must be applied in order to achieve the same current strength. The diffusion overvoltage can be reduced by continuous stirring or with rotating disk or cylinder electrodes.

The hydrogen and oxygen overvoltage do not remain constant on many metals. They sometimes even increase slightly after 60 minutes.

Cell resistance

The electrical resistance of an electrolysis cell hinders the flow of current ( Ohm's law ) and should therefore be minimized, otherwise energy is lost in the form of heat. The resistance of an electrolysis cell depends on the distance between the electrodes, the size of the electrode surface and the conductivity .

In general, the following applies to the calculation of the resistance of an electrolytic cell:

In distilled water, the conductivity is very low - the resistance is very high - and electrolysis is difficult.

Conductivity of some solutions
material Temperature ( K ) Conductivity (1 / (ohm cm))
Distilled water 273 10 −6 to 10 −7
1 M KCl (aq.) 293 0.102
0.1 M KCl (aq.) 293 0.017
1 M NaCl 291 0.074
1 M HCl (aq.) 298 0.332
1 M KOH (aq.) 291 0.184
4 M NaOH (aq.) 0.3500
1 MH 2 SO 4 (aq.) 0.3700

The conductivity of solutions of low concentrations can be calculated using the specific electrolytic conductivity or the equivalent conductivity of the ions. The conductivity of solutions of very high concentration must be determined experimentally. Although the conductivity of strong acids is higher than in basic solutions of the same concentration, many electrolyses - due to the anodic dissolution processes or the delayed oxygen formation or halogen oxidation in the acidic range - are mainly carried out in a basic medium.

Current density

In order to increase the economic efficiency of electrolytic processes, the processes should be carried out at the highest possible current densities. This can be achieved by increasing the conductivity by adding salt or by increasing the temperature (the specific conductivity increases by about 1–2% for every degree increase in temperature). The current density is often limited by the diffusion limit current. From knowledge of the diffusion limit current, dimensionless key figures can be determined in order to be able to calculate the conversion even for larger systems. There is an imputed optimal current density for every electrolysis, it is mostly not the maximum current density.

In order to obtain the cleanest and most compact metal deposits possible, a low current density should be used. This is especially important for gold, silver and copper covers. Metal deposits at high current densities form so-called spikes, poles, trees and these can lead to short circuits.

Often - especially in organic chemistry - thermal processes are superior to electrolytic processes due to the higher turnover per unit of time.

Ion migration speeds

During electrolysis, cations can be reduced on the cathode and anions can be oxidized on the anode. Since changes in charge due to reduction or oxidation occur close to the electrode, the charge difference in the electrode space must be compensated for by migration processes. Cations and anions must be present in identical concentrations in the electrode compartment; there must be no excess of positive or negative ions. The balance of ions in an electrolytic cell is brought about by ion migration. The speed of migration depends on the applied cell voltage and the type of ions. The loss of cations in front of the cathode can be compensated for by the migration of excess cations from the anode compartment or, conversely, by excess anions from the cathode compartment. As a rule, a compromise arises between the two directions of migration. The migration speeds can be calculated from the limit conductivities of the ion types. The change in the ion composition can be determined directly with the transfer number.

There are ions like H + or OH - that migrate very quickly in an electrolyte solution. Due to the different migration speeds, types of ions can accumulate in the half-cells of the electrolysis cell during electrolysis.

With a temperature increase of 1 ° C, the conductivity increases by approx. 1–2.5%. The increase in the migration speed could be justified with a lower viscosity of the solvation shell around the ions or even with a decrease in the solvation shell around the ions.

To link the parameters of migration speed , ion mobility (which is not a speed!), Electric field strength , equivalent conductivity / limiting conductivity (lambda) and the partial current brought about by ions in the electric field, see:

such as:

Examples

Electrolysis of water

The electrolysis of water breaks it down into the elements oxygen and hydrogen . Like all electrolyses, it consists of two partial reactions that take place on the two electrodes (cathode and anode compartments). The overall reaction scheme for this redox reaction is:

Electric current splits water into hydrogen and oxygen .

The electrodes are immersed in water, which is made more conductive by adding acid or lye . The partial reactions are

Cathode compartment : 2 H 3 O + + 2 e - → H 2 + 2 H 2 O (for acidic solutions) or: 2 H 2 O + 2 e - → H 2 + 2 OH - (for basic solutions)

Anode compartment: 6 H 2 O → O 2 + 4 H 3 O + + 4 e - (for acidic solutions) or: 4 OH - → O 2 + 2 H 2 O + 4 e - (for basic solutions)

As a demonstration experiment, this reaction can be carried out in Hofmann's water decomposition apparatus .

Water electrolysis can become more important as a storable energy carrier for the production of hydrogen. The energetic efficiency of the electrolysis of water is over 70%.

Electrolysis of zinc iodide

The electrolysis of zinc iodide breaks it down into the elements zinc and iodine . Like all electrolyses, this also consists of two partial reactions that take place on the two electrodes (cathode and anode compartment). The overall reaction scheme for this redox reaction is:

The zinc iodide was split into zinc and iodine by an electric current .

The reactions at the individual electrode spaces are:

Cathode reaction : Zn 2+ + 2 e - → Zn

Anode reaction: 2 I - → I 2 + 2 e -

The energy supply causes the individual ions to move towards the electrodes. The zinc cations migrate to the cathode, two electrons are absorbed by the zinc cations (reduction) and elemental zinc is formed. The iodine anions migrate to the anode and are oxidized to elemental iodine.

Applications

Substance extraction

A silver - Crystal electrodeposited with clearly visible dendritic structures.

The metals aluminum and magnesium are produced electrolytically with the help of fused- salt electrolysis . Copper , silver and gold are also obtained electrochemically , as well as zinc and nickel to a large extent . Other alkali metals and most alkaline earth metals are also obtained by fused-salt electrolysis.

Both during this process and during electrolysis in aqueous media, depending on the starting material, the halogens fluorine , bromine and chlorine are released, which are used on a large scale for further syntheses .

In the chlor-alkali electrolysis , chlorine, hydrogen and caustic soda are produced from rock salt .

Electroplating

Electrolytic metal deposition is one of the most important applications, either for the production of metallic coatings in electroplating ( electroplating , chrome plating, etc.) or for the production and reinforcement of conductor tracks in circuit board production .

Electrolytic refining

The electrolytic refining is a process for purifying metals. The cleaning is achieved in that an anode is released from a raw metal by electrolysis and is selectively deposited as pure metal on a cathode. Impurities remain dissolved in the electrolyte or precipitate as anode sludge . The anode sludge and electrolyte are processed because of their valuable components. Electrolytic refining is used in particular for the purification of copper , nickel , silver and lead .

During the electrolytic refining of copper , electrolytic copper with a purity of> 99.5% is obtained and is mainly used for electrical conductors . In copper refining, the cell voltage is a few tenths of a volt (mainly caused by overvoltages and cell resistance), the current density is in the range from 150 to 240 A / m 2 . The anode sludge contains the precious metals gold and silver in particular, but also selenium and antimony . The less noble metals, such as iron, nickel, cobalt and zinc, remain in the electrolyte.

In electrolytic lead refining , the refining of raw lead is used to separate arsenic, antimony and bismuth.

Kolbe electrolysis

The Kolbe electrolysis is the oldest example of an organic electrochemical reaction . During this electrolysis, two carboxylic acid molecules are coupled with the elimination of CO 2 .

Other uses of electrolysis

  • Analytical chemistry : With voltammetry and polarography , information about the chemical composition of the electrolyte is obtained by measuring the electrolysis current as a function of the voltage. In electrogravimetry and coulometry , the conversion of electrolytes by electrical current is used to obtain information about the metal content of a sample.
  • Wastewater treatment : In addition to hydroxide precipitation and the purification of wastewater with ion exchangers, electrochemical cleaning methods are used to purify contaminated wastewater from the metalworking industry, electroplating, dye and pharmaceutical industries. At the anode, cyanide salts and organic compounds are rendered harmless by oxidation. At the cathode z. B. lead, arsenic and copper removed by reduction, chromate is reduced to Cr 3+ .
  • Electrochemical Ablation : Electrochemical Ablation (ECM) is also called electrochemical metalworking. The workpiece is connected as an anode and the metal then dissolves due to its close proximity to the cathode. The shape of the cathode can influence the detachment at the anode. Suitable metals are aluminum, cobalt, molybdenum, nickel, titanium, tungsten, steel and iron alloys. Sodium nitrate or sodium hydroxide is used as the electrolyte .
  • Isotope separation : Natural water contains some deuterium . Since deuterium reacts much more slowly than hydrogen at the cathode to form the mixed gas molecule deuterium hydrogen, deuterium can be electrolytically enriched.
  • Charging batteries : When batteries are discharged, chemical energy is converted into usable electrical energy. If you recharge a battery, the charging voltage forces a reversal of the redox reaction that takes place during discharge. Thus, the charging process is electrolysis according to the definition given above. However, this naming is unusual.

economy

According to the Federal Statistical Office, the following quantities of metals or chemicals were produced in Germany in 2007.

material Production volume t (m 3 ) / year Sales value € million
Caustic soda (aqueous) 4,316,903 501
Chlorine gas 5,082,913 421
Potassium hydroxide (aqueous) 177.506 52
Aluminum (unalloyed) 279,660 529
Aluminum (alloyed) 1,033,860 1,397
Gold (as a semi-finished product) 91 901
Silver (as a semi-finished product) 2,635 455
Copper (refined) 553,300 1,629
Zinc (raw, refined) 264,843 654

In the USA, the electrolysis products manufactured are 2–3 times higher. About 5% of the total electricity produced there is required for electrolysis. The total amount of aluminum produced by electrolysis worldwide is around 50 million tons annually.

literature

  • Handbook of Experimental Chemistry. Upper secondary level. Volume 6: Electrochemistry. Aulis Verlag Deubner, 1994, ISBN 3-7614-1630-X .
  • Ullmann Encyclopedia of Industrial Chemistry. 3. Edition. Volume 6, 1955, pp. 253-304. (4th edition, volume 3, pp. 262–298, 5th edition, volume A9, p. 220 ff.)
  • Gerd Wedler: Textbook of physical chemistry. Verlag Chemie, 1982, ISBN 3-527-25880-9 , pp. 172-212, pp. 405-445, pp. 821-836.
  • Udo R. Kunze: Basics of quantitative analysis. Georg Thieme Verlag, 1980, ISBN 3-13-585801-4 , pp. 169-171.
  • Carl H. Hamann, Wolf Vielstich: Electrochemistry. 4th edition. WILEY-VCH Verlag, Weinheim 2005, ISBN 3-527-31068-1 .
  • Bernd Speiser: Electroanalytical Methods I: Electrode Reactions and Chronoamperometry . In: Chemistry in Our Time . tape 15 , no. 1 , February 1981, p. 21-26 , doi : 10.1002 / ciuz.19810150105 .
  • Wolfgang-Dieter Luz, Eberhard Zirngiebl: The future of electrochemistry . In: Chemistry in Our Time . tape 23 , no. 5 , October 1, 1989, pp. 151–160 , doi : 10.1002 / ciuz.19890230503 .

Web links

Commons : Electrolysis  - Album with pictures, videos and audio files
Wiktionary: Electrolysis  - explanations of meanings, word origins, synonyms, translations

Individual evidence

  1. ^ Schmidt-Walter: The alkaline fuel cell. 2005, p. 7.
  2. Rule of thumb: “Anode” and “Oxidation” begin with the vowels A and O , “Cathode” and “Reduction” with the consonants K and R, respectively .
  3. M. Binnewies, M. Jäckel, H. Willner, G. Rayner-Canham: Allgemeine und Anorganische Chemie . Ed .: Michael Binnewies. 2nd Edition. Spektrum Akademischer Verlag, Heidelberg 2011, ISBN 978-3-8274-2533-1 , p. 267 .
  4. ^ Ullmann Encyclopedia of Technical Chemistry. 3. Edition. Volume 6, p. 444.
  5. ^ A. Hickling, S. Hill: Oxygen Overvoltage. The influence of electrode material, current density and time in aqueous solutions. In: Discussions of the Faraday Society. No. 1, 1947.
  6. a b A. Hickling, FW Salt: Studies In Hydrogen Overvoltage At High Current Densities. Part I. In: Transactions of the Faraday Society. Volume 36, 1940, p. 1226.
  7. a b c d e f Gerd Wedler: Textbook of physical chemistry. 3. Edition. VCH, Weinheim 1987, p. 183.
  8. ^ Ullmann Encyclopedia of Technical Chemistry. 3. Edition. Volume 6, p. 450.
  9. ^ Ullmann Encyclopedia of Technical Chemistry. 4th edition. Volume 3, pp. 269-271, 281.
  10. ^ Ullmann Encyclopedia of Technical Chemistry. 3. Edition. Volume 6, p. 474.
  11. Philipp Lenard: About the conduction of electricity through free electrons and carriers. III: Migration speed of power-driven particles in friction media . With contributions by W. Weick and Hans Ferd. Mayer. In: Annals of Physics . tape 366 , no. 8 (episode 4 volume 61), 1920, p. 665–741 , here p. 718 , doi : 10.1002 / andp.19203660802 ( archive.org ).
  12. Alexander Stubinitzky: Eco-efficiency analysis of technical paths for the regenerative provision of hydrogen as fuel . In: progress reports VDI . 6: Energy technology , no. 588 . VDI Verlag, 2009, ISBN 978-3-18-358806-0 , ISSN  0178-9414 .
  13. ^ Arnold F. Holleman, Nils Wiberg: Textbook of Inorganic Chemistry. 102nd edition. de Gruyter, Berlin 2007, p. 1436 f.
  14. ^ Ullmann Encyclopedia of Technical Chemistry. 4th edition. Volume 3, p. 302.
  15. ^ Ullmann Encyclopedia of Technical Chemistry. 4th edition. Volume 3, p. 299.
  16. ^ Federal Statistical Office, Manufacturing Industry. Year 2007. Fachserie 4, Reihe 3.1, published: May 6, 2008.
  17. Kirk Othmer: Encyclopedia of Technology. 5th edition. 2005, Volume 9, p. 619.
  18. Primary Aluminum Production. The International Aluminum Institute, May 20, 2015, accessed June 9, 2015 .
  19. ^ Metallurgical production of aluminum by the most important countries in 2014 (in 1,000 tons). Statista, accessed June 9, 2015 .