Aluminium

Aluminium (aluminum in American and Canadian English) is a chemical element with the symbol Al and atomic number 13. It is a silvery-white, soft, non-magnetic and ductile metal in the boron group. By mass, aluminium makes up about 8% of the Earth's crust, where it is the third most abundant element (after oxygen and silicon) and also the most abundant metal. Occurrence of aluminium decreases in the Earth's mantle below, however. The chief ore of aluminium is bauxite. Aluminium metal is highly reactive, such that native specimens are rare and limited to extreme reducing environments. Instead, it is found combined in over 270 different minerals.[7]

For the alloys of aluminium, see aluminium alloy.

Aluminium, ₁₃Al

Aluminium
Pronunciation	aluminium: /ˌæljʊˈmɪniəm/ (listen) (AL-yuu-MIN-ee-əm) aluminum: /əˈljuːmɪnəm/ (listen) (ə-LEW-min-əm)
Alternative name	aluminum (U.S., Canada)
Appearance	silvery gray metallic
Standard atomic weight Ar, std(Al)	26.9815384(3)[1]
Aluminium in the periodic table
Hydrogen Helium Lithium Beryllium Boron Carbon Nitrogen Oxygen Fluorine Neon Sodium Magnesium Aluminium Silicon Phosphorus Sulfur Chlorine Argon Potassium Calcium Scandium Titanium Vanadium Chromium Manganese Iron Cobalt Nickel Copper Zinc Gallium Germanium Arsenic Selenium Bromine Krypton Rubidium Strontium Yttrium Zirconium Niobium Molybdenum Technetium Ruthenium Rhodium Palladium Silver Cadmium Indium Tin Antimony Tellurium Iodine Xenon Caesium Barium Lanthanum Cerium Praseodymium Neodymium Promethium Samarium Europium Gadolinium Terbium Dysprosium Holmium Erbium Thulium Ytterbium Lutetium Hafnium Tantalum Tungsten Rhenium Osmium Iridium Platinum Gold Mercury (element) Thallium Lead Bismuth Polonium Astatine Radon Francium Radium Actinium Thorium Protactinium Uranium Neptunium Plutonium Americium Curium Berkelium Californium Einsteinium Fermium Mendelevium Nobelium Lawrencium Rutherfordium Dubnium Seaborgium Bohrium Hassium Meitnerium Darmstadtium Roentgenium Copernicium Nihonium Flerovium Moscovium Livermorium Tennessine Oganesson B ↑ Al ↓ Ga magnesium ← aluminium → silicon
Atomic number (Z)	13
Group	group 13 (boron group)
Period	period 3
Block	p-block
Element category	Post-transition metal, [2][lower-alpha 1] sometimes considered a metalloid
Electron configuration	[Ne] 3s^2 3p^1
Electrons per shell	2, 8, 3
Physical properties
Phase at STP	solid
Melting point	933.47 K ​(660.32 °C, ​1220.58 °F)
Boiling point	2743 K ​(2470 °C, ​4478 °F)
Density (near r.t.)	2.70 g/cm^3
when liquid (at m.p.)	2.375 g/cm^3
Heat of fusion	10.71 kJ/mol
Heat of vaporization	284 kJ/mol
Molar heat capacity	24.20 J/(mol·K)
Vapor pressure P (Pa) 1 10 100 1 k 10 k 100 k at T (K) 1482 1632 1817 2054 2364 2790
Atomic properties
Oxidation states	−2, −1, +1,[4] +2,[5] +3 (an amphoteric oxide)
Electronegativity	Pauling scale: 1.61
Ionization energies	1st: 577.5 kJ/mol 2nd: 1816.7 kJ/mol 3rd: 2744.8 kJ/mol (more)
Atomic radius	empirical: 143 pm
Covalent radius	121±4 pm
Van der Waals radius	184 pm
Color lines in a spectral range Spectral lines of aluminium
Other properties
Natural occurrence	primordial
Crystal structure	​face-centered cubic (fcc)
Speed of sound thin rod	(rolled) 5000 m/s (at r.t.)
Thermal expansion	23.1 µm/(m·K) (at 25 °C)
Thermal conductivity	237 W/(m·K)
Electrical resistivity	26.5 nΩ·m (at 20 °C)
Magnetic ordering	paramagnetic[6]
Magnetic susceptibility	+16.5·10^−6 cm^3/mol
Young's modulus	70 GPa
Shear modulus	26 GPa
Bulk modulus	76 GPa
Poisson ratio	0.35
Mohs hardness	2.75
Vickers hardness	160–350 MPa
Brinell hardness	160–550 MPa
CAS Number	7429-90-5
History
Naming	after alumina (aluminium oxide), itself named after mineral alum
Prediction	Antoine Lavoisier (1782)
Discovery	Hans Christian Ørsted (1824)
First isolation	Friedrich Wöhler (1827, 1845)
Named by	Humphry Davy (1808, 1812)
Main isotopes of aluminium
Iso­tope Abun­dance Half-life (t1/2) Decay mode Pro­duct ^26Al trace 7.17×10^5 y β^+ ^26Mg ε ^26Mg γ – ^27Al 100% stable

Aluminium is remarkable for its low density and its ability to resist corrosion through the phenomenon of passivation. Aluminium and its alloys are vital to the aerospace industry[8] and important in transportation and building industries, such as building facades and window frames.[9] The oxides and sulfates are the most useful compounds of aluminium.[8]

Despite its prevalence in the environment, no known form of life uses aluminium salts metabolically, but aluminium is well tolerated by plants and animals.[10] Because of these salts' abundance, the potential for a biological role for them is of continuing interest, and studies continue.

Physical characteristics

Isotopes

Main article: Isotopes of aluminium

Of aluminium isotopes, only ²⁷
Al
is stable. This is consistent with aluminium having an odd atomic number.[lower-alpha 2] It is the only primordial aluminium isotope, i.e. the only one that has existed on Earth in its current form since the creation of the planet. Nearly all aluminium on Earth is present as this isotope, which makes it a mononuclidic element and means that its standard atomic weight is the same as that of the isotope. The standard atomic weight of aluminium is low in comparison with many other metals,[lower-alpha 3] which has consequences for the element's properties (see below). This makes aluminium very useful in nuclear magnetic resonance (NMR), as its single stable isotope has a high NMR sensitivity.[12]

All other isotopes of aluminium are radioactive. The most stable of these is ²⁶Al: while it was present along with stable ²⁷Al in the interstellar medium from which the Solar System formed, having been produced by stellar nucleosynthesis as well, its half-life is only 717,000 years and therefore it could not have survived since the formation of the planet. However, minute traces of ²⁶Al are produced from argon in the atmosphere by spallation caused by cosmic ray protons. The ratio of ²⁶Al to ¹⁰Be has been used for radiodating of geological processes over 10⁵ to 10⁶ year time scales, in particular transport, deposition, sediment storage, burial times, and erosion.[13] Most meteorite scientists believe that the energy released by the decay of ²⁶Al was responsible for the melting and differentiation of some asteroids after their formation 4.55 billion years ago.[14]

The remaining isotopes of aluminium, with mass numbers ranging from 22 to 43, all have half-lives well under an hour. Three metastable states are known, all with half-lives under a minute.[11]

Electron shell

An aluminium atom has 13 electrons, arranged in an electron configuration of [Ne] 3s² 3p¹,[15] with three electrons beyond a stable noble gas configuration. Accordingly, the combined first three ionization energies of aluminium are far lower than the fourth ionization energy alone.[16] Such an electron configuration is shared with the other well-characterized members of its group, boron, gallium, indium, and thallium; it is also expected for nihonium. Aluminium can relatively easily surrender its three outermost electrons in many chemical reactions (see below). The electronegativity of aluminium is 1.61 (Pauling scale).[17]

High-resolution STEM-HAADF micrograph of Al atoms viewed along the [001] zone axis.

A free aluminium atom has a radius of 143 pm.[18] With the three outermost electrons removed, the radius shrinks to 39 pm for a 4-coordinated atom or 53.5 pm for a 6-coordinated atom.[18] At standard temperature and pressure, aluminium atoms (when not affected by atoms of other elements) form a face-centered cubic crystal system bound by metallic bonding provided by atoms' outermost electrons; hence aluminium (at these conditions) is a metal.[19] This crystal system is shared by many other metals, such as lead and copper; the size of a unit cell of aluminium is comparable to that of those other metals.[19] It is however not shared by the other members of its group; boron has ionization energies too high to allow metallization, thallium has a hexagonal close-packed structure, and gallium and indium have unusual structures that are not close-packed like those of aluminium and thallium. Since few electrons are available for metallic bonding, aluminium metal is soft with a low melting point and low electrical resistivity, as is common for post-transition metals.[20]

Bulk

Aluminium metal has an appearance ranging from silvery white to dull gray, depending on the surface roughness. A fresh film of aluminium serves as a good reflector (approximately 92%) of visible light and an excellent reflector (as much as 98%) of medium and far infrared radiation.

The density of aluminium is 2.70 g/cm³, about 1/3 that of steel, much lower than other commonly encountered metals, making aluminium parts easily identifiable through their lightness.[21] Aluminium's low density compared to most other metals arises from the fact that its nuclei are much lighter, while difference in the unit cell size does not compensate for this difference. The only lighter metals are the metals of groups 1 and 2, which apart from beryllium and magnesium are too reactive for structural use (and beryllium is very toxic).[22] Aluminium is not as strong or stiff as steel, but the low density makes up for this in the aerospace industry and for many other applications where light weight and relatively high strength are crucial.

Pure aluminium is quite soft and lacking in strength. In most applications various aluminium alloys are used instead because of their higher strength and hardness. The yield strength of pure aluminium is 7–11 MPa, while aluminium alloys have yield strengths ranging from 200 MPa to 600 MPa.[23] Aluminium is ductile, with a percent elongation of 50-70%,[24] and malleable allowing it to be easily drawn and extruded. It is also easily machined, and the low melting temperature of 660 °C allows for easy casting.

Aluminium is an excellent thermal and electrical conductor, having 59% the conductivity of copper, both thermal and electrical, while having only 30% of copper's density. Aluminium is capable of superconductivity, with a superconducting critical temperature of 1.2 kelvin and a critical magnetic field of about 100 gauss (10 milliteslas).[25] It is paramagnetic and thus essentially unaffected by static magnetic fields. The high electrical conductivity, however, means that it is strongly affected by alternating magnetic fields through the induction of eddy currents.

Chemistry

Main article: Compounds of aluminium

Aluminium combines characteristics of pre- and post-transition metals. Since it has few available electrons for metallic bonding, like its heavier group 13 congeners, it has the characteristic physical properties of a post-transition metal, with longer-than-expected interatomic distances.[20] Furthermore, as Al³⁺ is a small and highly charged cation, it is strongly polarizing and aluminium compounds tend towards covalency;[26] this behaviour is similar to that of beryllium (Be²⁺), and the two display an example of a diagonal relationship.[27]

The underlying core under aluminium's valence shell is that of the preceding noble gas, whereas those of its heavier congeners gallium and indium, thallium, and nihonium also include a filled d-subshell and in some cases a filled f-subshell. Hence, the inner electrons of aluminium shield the valence electrons almost completely, unlike those of aluminium's heavier congeners. As such, aluminium is the most electropositive metal in its group, and its hydroxide is in fact more basic than that of gallium.[26] In fact, aluminium's electropositive behavior, high affinity for oxygen, and highly negative standard electrode potential are all more similar to those of scandium, yttrium, lanthanum, and actinium, which like aluminium have three valence electrons outside a noble gas core.[20] Aluminium also bears minor similarities to the metalloid boron in the same group: AlX₃ compounds are valence isoelectronic to BX₃ compounds (they have the same valence electronic structure), and both behave as Lewis acids and readily form adducts.[28] Additionally, one of the main motifs of boron chemistry is regular icosahedral structures, and aluminium forms an important part of many icosahedral quasicrystal alloys, including the Al–Zn–Mg class.[29]

Aluminium has a high chemical affinity to oxygen, which renders it suitable for use as a reducing agent in the thermite reaction. A fine powder of aluminium metal reacts explosively on contact with liquid oxygen; under normal conditions, however, aluminium forms a thin oxide layer (~ 5 nm at room temperature)[30] that protects the metal from further corrosion by oxygen, water, or dilute acid, a process termed passivation.[26][31] Because of its general resistance to corrosion, aluminium is one of the few metals that retains silvery reflectance in finely powdered form, making it an important component of silver-colored paints.[32] Aluminium is not attacked by oxidizing acids because of its passivation. This allows aluminium to be used to store reagents such as nitric acid, concentrated sulfuric acid, and some organic acids.[10]

In hot concentrated hydrochloric acid, aluminium reacts with water with evolution of hydrogen, and in aqueous sodium hydroxide or potassium hydroxide at room temperature to form aluminates—protective passivation under these conditions is negligible.[33] Aqua regia also dissolves aluminium.[10] Aluminium is corroded by dissolved chlorides, such as common sodium chloride, which is why household plumbing is never made from aluminium.[33] The oxide layer on aluminium is also destroyed by contact with mercury due to amalgamation or with salts of some electropositive metals.[26] As such, the strongest aluminium alloys are less corrosion-resistant due to galvanic reactions with alloyed copper,[23] and aluminium's corrosion resistance is greatly reduced by aqueous salts, particularly in the presence of dissimilar metals.[20]

Aluminium reacts with most nonmetals upon heating, forming compounds such as aluminium nitride (AlN), aluminium sulfide (Al₂S₃), and the aluminium halides (AlX₃). It also forms a wide range of intermetallic compounds involving metals from every group on the periodic table.[26]

Inorganic compounds

The vast majority of compounds, including all aluminium-containing minerals and all commercially significant aluminium compounds, feature aluminium in the oxidation state 3+. The coordination number of such compounds varies, but generally Al³⁺ is either six- or four-coordinate. Almost all compounds of aluminium(III) are colorless.[26]

Aluminium hydrolysis as a function of pH. Coordinated water molecules are omitted. (Data from Baes and Mesmer)[34]

In aqueous solution, Al³⁺ exists as the hexaaqua cation [Al(H₂O)₆]³⁺, which has an approximate pK_a of 10⁻⁵.[12] Such solutions are acidic as this cation can act as a proton donor and progressively hydrolyse until a precipitate of aluminium hydroxide, Al(OH)₃, forms. This is useful for clarification of water, as the precipitate nucleates on suspended particles in the water, hence removing them. Increasing the pH even further leads to the hydroxide dissolving again as aluminate, [Al(H₂O)₂(OH)₄]⁻, is formed.

Aluminium hydroxide forms both salts and aluminates and dissolves in acid and alkali, as well as on fusion with acidic and basic oxides.[26] This behaviour of Al(OH)₃ is termed amphoterism, and is characteristic of weakly basic cations that form insoluble hydroxides and whose hydrated species can also donate their protons. One effect of this is that aluminium salts with weak acids are hydrolysed in water to the aquated hydroxide and the corresponding nonmetal hydride: for example, aluminium sulfide yields hydrogen sulfide. However, some salts like aluminium carbonate exist in aqueous solution but are unstable as such; and only incomplete hydrolysis takes place for salts with strong acids, such as the halides, nitrate, and sulfate. For similar reasons, anhydrous aluminium salts cannot be made by heating their "hydrates": hydrated aluminium chloride is in fact not AlCl₃·6H₂O but [Al(H₂O)₆]Cl₃, and the Al–O bonds are so strong that heating is not sufficient to break them and form Al–Cl bonds instead:[26]

2[Al(H₂O)₆]Cl₃ heat→  Al₂O₃ + 6 HCl + 9 H₂O

All four trihalides are well known. Unlike the structures of the three heavier trihalides, aluminium fluoride (AlF₃) features six-coordinate aluminium, which explains its involatility and insolubility as well as high heat of formation. Each aluminium atom is surrounded by six fluorine atoms in a distorted octahedral arrangement, with each fluorine atom being shared between the corners of two octahedra. Such {AlF₆} units also exist in complex fluorides such as cryolite, Na₃AlF₆.[lower-alpha 4] AlF₃ melts at 1,290 °C (2,354 °F) and is made by reaction of aluminium oxide with hydrogen fluoride gas at 700 °C (1,292 °F).[35]

With heavier halides, the coordination numbers are lower. The other trihalides are dimeric or polymeric with tetrahedral four-coordinate aluminium centers. Aluminium trichloride (AlCl₃) has a layered polymeric structure below its melting point of 192.4 °C (378 °F) but transforms on melting to Al₂Cl₆ dimers. At higher temperatures those increasingly dissociate into trigonal planar AlCl₃ monomers similar to the structure of BCl₃. Aluminium tribromide and aluminium triiodide form Al₂X₆ dimers in all three phases and hence do not show such significant changes of properties upon phase change.[35] These materials are prepared by treating aluminium metal with the halogen. The aluminium trihalides form many addition compounds or complexes; their Lewis acidic nature makes them useful as catalysts for the Friedel–Crafts reactions. Aluminium trichloride has major industrial uses involving this reaction, such as in the manufacture of anthraquinones and styrene; it is also often used as the precursor for many other aluminium compounds and as a reagent for converting nonmetal fluorides into the corresponding chlorides (a transhalogenation reaction).[35]

Aluminium forms one stable oxide with the chemical formula Al₂O₃, commonly called alumina.[36] It can be found in nature in the mineral corundum, α-alumina;[37] there is also a γ-alumina phase.[12] Its crystalline form, corundum, is very hard (Mohs hardness 9), has a high melting point of 2,045 °C (3,713 °F), has very low volatility, is chemically inert, and a good electrical insulator, it is often used in abrasives (such as toothpaste), as a refractory material, and in ceramics, as well as being the starting material for the electrolytic production of aluminium metal. Sapphire and ruby are impure corundum contaminated with trace amounts of other metals.[12] The two main oxide-hydroxides, AlO(OH), are boehmite and diaspore. There are three main trihydroxides: bayerite, gibbsite, and nordstrandite, which differ in their crystalline structure (polymorphs). Many other intermediate and related structures are also known.[12] Most are produced from ores by a variety of wet processes using acid and base. Heating the hydroxides leads to formation of corundum. These materials are of central importance to the production of aluminium and are themselves extremely useful. Some mixed oxide phases are also very useful, such as spinel (MgAl₂O₄), Na-β-alumina (NaAl₁₁O₁₇), and tricalcium aluminate (Ca₃Al₂O₆, an important mineral phase in Portland cement).[12]

The only stable chalcogenides under normal conditions are aluminium sulfide (Al₂S₃), selenide (Al₂Se₃), and telluride (Al₂Te₃). All three are prepared by direct reaction of their elements at about 1,000 °C (1,832 °F) and quickly hydrolyse completely in water to yield aluminium hydroxide and the respective hydrogen chalcogenide. As aluminium is a small atom relative to these chalcogens, these have four-coordinate tetrahedral aluminium with various polymorphs having structures related to wurtzite, with two-thirds of the possible metal sites occupied either in an orderly (α) or random (β) fashion; the sulfide also has a γ form related to γ-alumina, and an unusual high-temperature hexagonal form where half the aluminium atoms have tetrahedral four-coordination and the other half have trigonal bipyramidal five-coordination.[38]

Four pnictides – aluminium nitride (AlN), aluminium phosphide (AlP), aluminium arsenide (AlAs), and aluminium antimonide (AlSb) – are known. They are all III-V semiconductors isoelectronic to silicon and germanium, all of which but AlN have the zinc blende structure. All four can be made by high-temperature (and possibly high-pressure) direct reaction of their component elements.[38]

Rarer oxidation states

See also: Aluminium(I)

Although the great majority of aluminium compounds feature Al³⁺ centers, compounds with lower oxidation states are known and are sometimes of significance as precursors to the Al³⁺ species.

AlF, AlCl, AlBr, and AlI exist in the gaseous phase when the respective trihalide is heated with aluminium, and at cryogenic temperatures. Their instability in the condensed phase is due to their ready disproportionation to aluminium and the respective trihalide: the reverse reaction is favored at high temperature (although even then they are still short-lived), explaining why AlF₃ is more volatile when heated in the presence of aluminium metal, as is aluminium metal when heated in the presence of AlCl₃.[35] A stable derivative of aluminium monoiodide is the cyclic adduct formed with triethylamine, Al₄I₄(NEt₃)₄. Also of theoretical interest but only of fleeting existence are Al₂O and Al₂S. Al₂O is made by heating the normal oxide, Al₂O₃, with silicon at 1,800 °C (3,272 °F) in a vacuum. Such materials quickly disproportionate to the starting materials.[39]

Very simple Al(II) compounds are invoked or observed in the reactions of Al metal with oxidants. For example, aluminium monoxide, AlO, has been detected in the gas phase after explosion[40] and in stellar absorption spectra.[41] More thoroughly investigated are compounds of the formula R₄Al₂ which contain an Al–Al bond and where R is a large organic ligand.[42]

Organoaluminium compounds and related hydrides

Structure of trimethylaluminium, a compound that features five-coordinate carbon.

A variety of compounds of empirical formula AlR₃ and AlR_1.5Cl_1.5 exist.[43] The aluminium trialkyls and triaryls are reactive, volatile, and colorless liquids or low-melting solids. They catch fire spontaneously in air and react with water, thus necessitating precautions when handling them. They often form dimers, unlike their boron analogues, but this tendency diminishes for branched-chain alkyls (e.g. Prⁱ, Buⁱ, Me₃CCH₂); for example, triisobutylaluminium exists as an equilibrium mixture of the monomer and dimer.[44][45] These dimers, such as trimethylaluminium (Al₂Me₆), usually feature tetrahedral Al centers formed by dimerization with some alkyl group bridging between both aluminium atoms. They are hard acids and react readily with ligands, forming adducts. In industry, they are mostly used in alkene insertion reactions, as discovered by Karl Ziegler, most importantly in "growth reactions" that form long-chain unbranched primary alkenes and alcohols, and in the low-pressure polymerization of ethene and propene. There are also some heterocyclic and cluster organoaluminium compounds involving Al–N bonds.[44]

The industrially most important aluminium hydride is lithium aluminium hydride (LiAlH₄), which is used in as a reducing agent in organic chemistry. It can be produced from lithium hydride and aluminium trichloride.[46] The simplest hydride, aluminium hydride or alane, is not as important. It is a polymer with the formula (AlH₃)_n, in contrast to the corresponding boron hydride that is a dimer with the formula (BH₃)₂.[46]

Natural occurrence

See also: List of countries by bauxite production

In space

Aluminium's per-particle abundance in the Solar System is 3.15 ppm (parts per million).[47][lower-alpha 5] It is the twelfth most abundant of all elements and third most abundant among the elements that have odd atomic numbers, after hydrogen and nitrogen.[47] The only stable isotope of aluminium, ²⁷Al, is the eighteenth most abundant nucleus in the Universe. It is created almost entirely after fusion of carbon in massive stars that will later become Type II supernovae: this fusion creates ²⁶Mg, which, upon capturing free protons and neutrons becomes aluminium. Some smaller quantities of ²⁷Al are created in hydrogen burning shells of evolved stars, where ²⁶Mg can capture free protons.[48] Essentially all aluminium now in existence is ²⁷Al; ²⁶Al was present in the early Solar System but is currently extinct. However, the trace quantities of ²⁶Al that do exist are the most common gamma ray emitter in the interstellar gas.[48]

On Earth

Bauxite, a major aluminium ore. The red-brown color is due to the presence of iron oxide minerals.

Overall, the Earth is about 1.59% aluminium by mass (seventh in abundance by mass).[49] Aluminium occurs in greater proportion in the Earth than in the Universe because aluminium easily forms the oxide and becomes bound into rocks and aluminium stays in the Earth's crust while less reactive metals sink to the core.[48] In the Earth's crust, aluminium is the most abundant (8.23% by mass[24]) metallic element and the third most abundant of all elements (after oxygen and silicon).[50] A large number of silicates in the Earth's crust contain aluminium.[51] In contrast, the Earth's mantle is only 2.38% aluminium by mass.[52] Aluminium also occurs in seawater at a concentration of 2 μg/kg.[24]

Because of its strong affinity for oxygen, aluminium is almost never found in the elemental state; instead it is found in oxides or silicates. Feldspars, the most common group of minerals in the Earth's crust, are aluminosilicates. Aluminium also occurs in the minerals beryl, cryolite, garnet, spinel, and turquoise.[53] Impurities in Al₂O₃, such as chromium and iron, yield the gemstones ruby and sapphire, respectively.[54] Native aluminium metal can only be found as a minor phase in low oxygen fugacity environments, such as the interiors of certain volcanoes.[55] Native aluminium has been reported in cold seeps in the northeastern continental slope of the South China Sea. It is possible that these deposits resulted from bacterial reduction of tetrahydroxoaluminate Al(OH)₄⁻.[56]

Although aluminium is a common and widespread element, not all aluminium minerals are economically viable sources of the metal. Almost all metallic aluminium is produced from the ore bauxite (AlO_x(OH)_3–2x). Bauxite occurs as a weathering product of low iron and silica bedrock in tropical climatic conditions.[57] In 2017, most bauxite was mined in Australia, China, Guinea, and India.[58]

History

Main article: History of aluminium

Friedrich Wöhler, the chemist who first thoroughly described metallic elemental aluminium

The history of aluminium has been shaped by usage of alum. The first written record of alum, made by Greek historian Herodotus, dates back to the 5th century BCE.[59] The ancients are known to have used alum as a dyeing mordant and for city defense.[59] After the Crusades, alum, an indispensable good in the European fabric industry,[60] was a subject of international commerce;[61] it was imported to Europe from the eastern Mediterranean until the mid-15th century.[62]

The nature of alum remained unknown. Around 1530, Swiss physician Paracelsus suggested alum was a salt of an earth of alum.[63] In 1595, German doctor and chemist Andreas Libavius experimentally confirmed this.[64] In 1722, German chemist Friedrich Hoffmann announced his belief that the base of alum was a distinct earth.[65] In 1754, German chemist Andreas Sigismund Marggraf synthesized alumina by boiling clay in sulfuric acid and subsequently adding potash.[65]

Attempts to produce aluminium metal date back to 1760.[66] The first successful attempt, however, was completed in 1824 by Danish physicist and chemist Hans Christian Ørsted. He reacted anhydrous aluminium chloride with potassium amalgam, yielding a lump of metal looking similar to tin.[67][68][69] He presented his results and demonstrated a sample of the new metal in 1825.[70][71] In 1827, German chemist Friedrich Wöhler repeated Ørsted's experiments but did not identify any aluminium.[72] (The reason for this inconsistency was only discovered in 1921.)[73] He conducted a similar experiment in the same year by mixing anhydrous aluminium chloride with potassium and produced a powder of aluminium.[69] In 1845, he was able to produce small pieces of the metal and described some physical properties of this metal.[73] For many years thereafter, Wöhler was credited as the discoverer of aluminium.[74]

The statue of Anteros in Piccadilly Circus, London, was made in 1893 and is one of the first statues cast in aluminium.

As Wöhler's method could not yield great quantities of aluminium, the metal remained rare; its cost exceeded that of gold.[72] The first industrial production of aluminium was established in 1856 by French chemist Henri Etienne Sainte-Claire Deville and companions.[75] Deville had discovered that aluminium trichloride could be reduced by sodium, which was more convenient and less expensive than potassium, which Wöhler had used.[76] Even then, aluminium was still not of great purity and produced aluminium differed in properties by sample.[77]

The first industrial large-scale production method was independently developed in 1886 by French engineer Paul Héroult and American engineer Charles Martin Hall; it is now known as the Hall–Héroult process.[78] The Hall–Héroult process converts alumina into the metal. Austrian chemist Carl Joseph Bayer discovered a way of purifying bauxite to yield alumina, now known as the Bayer process, in 1889.[79] Modern production of the aluminium metal is based on the Bayer and Hall–Héroult processes.[80]

Prices of aluminium dropped and aluminium became widely used in jewelry, everyday items, eyeglass frames, optical instruments, tableware, and foil in the 1890s and early 20th century. Aluminium's ability to form hard yet light alloys with other metals provided the metal many uses at the time.[81] During World War I, major governments demanded large shipments of aluminium for light strong airframes.[82]

By the mid-20th century, aluminium had become a part of everyday life and an essential component of housewares.[83] During the mid-20th century, aluminium emerged as a civil engineering material, with building applications in both basic construction and interior finish work,[84] and increasingly being used in military engineering, for both airplanes and land armor vehicle engines.[85] Earth's first artificial satellite, launched in 1957, consisted of two separate aluminium semi-spheres joined together and all subsequent space vehicles have used aluminium to some extent.[80] The aluminium can was invented in 1956 and employed as a storage for drinks in 1958.[86]

World production of aluminium since 1900

Throughout the 20th century, the production of aluminium rose rapidly: while the world production of aluminium in 1900 was 6,800 metric tons, the annual production first exceeded 100,000 metric tons in 1916; 1,000,000 tons in 1941; 10,000,000 tons in 1971.[87] In the 1970s, the increased demand for aluminium made it an exchange commodity; it entered the London Metal Exchange, the oldest industrial metal exchange in the world, in 1978.[80] The output continued to grow: the annual production of aluminium exceeded 50,000,000 metric tons in 2013.[87]

The real price for aluminium declined from $14,000 per metric ton in 1900 to $2,340 in 1948 (in 1998 United States dollars).[87] Extraction and processing costs were lowered over technological progress and the scale of the economies. However, the need to exploit lower-grade poorer quality deposits and the use of fast increasing input costs (above all, energy) increased the net cost of aluminium;[88] the real price began to grow in the 1970s with the rise of energy cost.[89] Production moved from the industrialized countries to countries where production was cheaper.[90] Production costs in the late 20th century changed because of advances in technology, lower energy prices, exchange rates of the United States dollar, and alumina prices.[91] The BRIC countries' combined share in primary production and primary consumption grew substantially in the first decade of the 21st century.[92] China is accumulating an especially large share of world's production thanks to abundance of resources, cheap energy, and governmental stimuli;[93] it also increased its consumption share from 2% in 1972 to 40% in 2010.[94] In the United States, Western Europe, and Japan, most aluminium was consumed in transportation, engineering, construction, and packaging.[95]

Etymology

Aluminium is named after alumina, a naturally occurring oxide of aluminium, and the name alumina comes from alum, the mineral from which it was collected. The word "alum" derives from the Latin word alumen, meaning "bitter salt".[96] The word alumen stems from the Proto-Indo-European root *alu- meaning "bitter" or "beer".[97]

1897 American advertisement featuring the aluminum spelling

Coinage

British chemist Humphry Davy, who performed a number of experiments aimed to isolate the metal, is credited as the person who named the element. The first named proposed for the metal to be isolated from alum was alumium, which Davy suggested in an 1808 article on his electrochemical research, published in Philosophical Transactions of the Royal Society.[98] This suggestion was criticized by contemporary chemists from France, Germany, and Sweden, who insisted the metal should be named for the oxide, alumina, from which it would be isolated.[99] A January 1811 summary of one of Davy's lectures at the Royal Society proposed the name aluminium [100]—this is the earliest known published writing to use either of the modern spellings. However, the following year, Davy published a chemistry textbook in which he settled on the spelling aluminum.[101] Both spellings have coexisted since; however, their usage has split by region: aluminum is in use in the United States and Canada while aluminium is in use elsewhere.[102]

Spelling

Davy's spelling aluminum is consistent with the Latin naming of metals, which end in -um, e.g. aurum (gold), argentum (silver), ferrum (iron),[103] naming newly discovered elements by replacing a -a or -ite suffix in the oxide's name with -um: lanthanum was named for its oxide lanthana, magnesium for magnesia, tantalum for tantalite, molybdenum for molybdenite (also known as molybdena), cerium for ceria, and thorium for thoria, respectively. As aluminium's oxide is called alumina, not aluminia, the -ium spelling does not follow this pattern. However, other newly discovered elements of the time had names with a -ium suffix, such as potassium, sodium, calcium, and strontium.

In 1812, British scientist Thomas Young[104] wrote an anonymous review of Davy's book, in which he proposed the name aluminium instead of aluminum, which he felt had a "less classical sound".[105] This name did catch on: while the -um spelling was occasionally used in Britain, the American scientific language used -ium from the start.[106] Most scientists used -ium throughout the world in the 19th century;[107] it still remains the standard in most other languages.[102] In 1828, American lexicographer Noah Webster used exclusively the aluminum spelling in his American Dictionary of the English Language.[108] In the 1830s, the -um spelling started to gain usage in the United States; by the 1860s, it had become the more common spelling there outside science.[106] In 1892, Hall used the -um spelling in his advertising handbill for his new electrolytic method of producing the metal, despite his constant use of the -ium spelling in all the patents he filed between 1886 and 1903. It was subsequently suggested this was a typo rather than intended.[102] By 1890, both spellings had been common in the U.S. overall, the -ium spelling being slightly more common; by 1895, the situation had reversed; by 1900, aluminum had become twice as common as aluminium; during the following decade, the -um spelling dominated American usage.[109] In 1925, the American Chemical Society adopted this spelling.[109]

The International Union of Pure and Applied Chemistry (IUPAC) adopted aluminium as the standard international name for the element in 1990.[110] In 1993, they recognized aluminum as an acceptable variant;[110] the most recent 2005 edition of the IUPAC nomenclature of inorganic chemistry acknowledges this spelling as well.[111] IUPAC official publications use the -ium spelling as primary but list both where appropriate.[lower-alpha 6]

Production and refinement

See also: List of countries by primary aluminium production

World's top producers of primary aluminium, 2016[113]

Country	Output (thousand tons)
China	31,873
Russia	3,561
Canada	3,208
India	2,896
United Arab Emirates	2,471
Australia	1,635
Norway	1,247
Bahrain	971
Saudi Arabia	869
United States	818
Brazil	793
South Africa	701
Iceland	700
World total	58,800

Aluminium production is highly energy-consuming, and so the producers tend to locate smelters in places where electric power is both plentiful and inexpensive.[114] As of 2012, the world's largest smelters of aluminium are located in China, Russia, Bahrain, United Arab Emirates, and South Africa.[115]

In 2016, China was the top producer of aluminium with a world share of fifty-five percent; the next largest producing countries were Russia, Canada, India, and the United Arab Emirates.[113]

According to the International Resource Panel's Metal Stocks in Society report, the global per capita stock of aluminium in use in society (i.e. in cars, buildings, electronics, etc.) is 80 kg (180 lb). Much of this is in more-developed countries (350–500 kg (770–1,100 lb) per capita) rather than less-developed countries (35 kg (77 lb) per capita).[116]

Bayer process

Main article: Bayer process

Bauxite is converted to aluminium oxide by the Bayer process. Bauxite is blended for uniform composition and then is ground. The resulting slurry is mixed with a hot solution of sodium hydroxide; the mixture is then treated in a digester vessel at a pressure well above atmospheric, dissolving the aluminium hydroxide in bauxite while converting impurities into relatively insoluble compounds:[117]

Al(OH)₃ + Na⁺ + OH⁻ → Na⁺ + [Al(OH)₄]⁻

After this reaction, the slurry is at a temperature above its atmospheric boiling point. It is cooled by removing steam as pressure is reduced. The bauxite residue is separated from the solution and discarded. The solution, free of solids, is seeded with small crystals of aluminium hydroxide; this causes decomposition of the [Al(OH)₄]⁻ ions to aluminium hydroxide. After about half of aluminium has precipitated, the mixture is sent to classifiers. Small crystals of aluminium hydroxide are collected to serve as seeding agents; coarse particles are converted to aluminium oxide by heating; excess solution is removed by evaporation, (if needed) purified, and recycled.[117]

Hall–Héroult process

Main articles: Hall–Héroult process and Aluminium smelting

The conversion of alumina to aluminium metal is achieved by the Hall–Héroult process. In this energy-intensive process, a solution of alumina in a molten (950 and 980 °C (1,740 and 1,800 °F)) mixture of cryolite (Na₃AlF₆) with calcium fluoride is electrolyzed to produce metallic aluminium. The liquid aluminium metal sinks to the bottom of the solution and is tapped off, and usually cast into large blocks called aluminium billets for further processing.[10]

Extrusion billets of aluminium

Anodes of the electrolysis cell are made of carbon—the most resistant material against fluoride corrosion—and either bake at the process or are prebaked. The former, also called Söderberg anodes, are less power-efficient and fumes released during baking are costly to collect, which is why they are being replaced by prebaked anodes even though they save the power, energy, and labor to prebake the cathodes. Carbon for anodes should be preferably pure so that neither aluminium nor the electrolyte is contaminated with ash. Despite carbon's resistivity against corrosion, it is still consumed at a rate of 0.4–0.5 kg per each kilogram of produced aluminium. Cathodes are made of anthracite; high purity for them is not required because impurities leach only very slowly. Cathode is consumed at a rate of 0.02–0.04 kg per each kilogram of produced aluminium. A cell is usually a terminated after 2–6 years following a failure of the cathode.[10]

The Hall–Heroult process produces aluminium with a purity of above 99%. Further purification can be done by the Hoopes process. This process involves the electrolysis of molten aluminium with a sodium, barium, and aluminium fluoride electrolyte. The resulting aluminium has a purity of 99.99%.[10][118]

Electric power represents about 20 to 40% of the cost of producing aluminium, depending on the location of the smelter. Aluminium production consumes roughly 5% of electricity generated in the United States.[110] Because of this, alternatives to the Hall–Héroult process have been researched, but none has turned out to be economically feasible.[10]

Common bins for recyclable waste along with a bin for unrecyclable waste. The bin with a yellow top is labeled "aluminum". Rhodes, Greece.

Recycling

Main article: Aluminium recycling

Recovery of the metal through recycling has become an important task of the aluminium industry. Recycling was a low-profile activity until the late 1960s, when the growing use of aluminium beverage cans brought it to public awareness.[119] Recycling involves melting the scrap, a process that requires only 5% of the energy used to produce aluminium from ore, though a significant part (up to 15% of the input material) is lost as dross (ash-like oxide).[120] An aluminium stack melter produces significantly less dross, with values reported below 1%.[121]

White dross from primary aluminium production and from secondary recycling operations still contains useful quantities of aluminium that can be extracted industrially. The process produces aluminium billets, together with a highly complex waste material. This waste is difficult to manage. It reacts with water, releasing a mixture of gases (including, among others, hydrogen, acetylene, and ammonia), which spontaneously ignites on contact with air;[122] contact with damp air results in the release of copious quantities of ammonia gas. Despite these difficulties, the waste is used as a filler in asphalt and concrete.[123]

Applications

Aluminium-bodied Austin A40 Sports (c. 1951)

Metal

See also: Aluminium alloy

The global production of aluminium in 2016 was 58.8 million metric tons. It exceeded that of any other metal except iron (1,231 million metric tons).[113][124]

Aluminium is almost always alloyed, which markedly improves its mechanical properties, especially when tempered. For example, the common aluminium foils and beverage cans are alloys of 92% to 99% aluminium.[125] The main alloying agents are copper, zinc, magnesium, manganese, and silicon (e.g., duralumin) with the levels of other metals in a few percent by weight.[126]

Aluminium can

The major uses for aluminium metal are in:[127]

• Transportation (automobiles, aircraft, trucks, railway cars, marine vessels, bicycles, spacecraft, etc.). Aluminium is used because of its low density;
• Packaging (cans, foil, frame etc.). Aluminium is used because it is non-toxic , non-adsorptive, and splinter-proof;
• Building and construction (windows, doors, siding, building wire, sheathing, roofing, etc.). Since steel is cheaper, aluminium is used when lightness, corrosion resistance, or engineering features are important;
• Electricity-related uses (conductor alloys, motors and generators, transformers, capacitors, etc.). Aluminium is used because it is relatively cheap, highly conductive, has adequate mechanical strength and low density, and resists corrosion;
• A wide range of household items, from cooking utensils to furniture. Low density, good appearance, ease of fabrication, and durability are the key factors of aluminium usage;
• Machinery and equipment (processing equipment, pipes, tools). Aluminium is used because of its corrosion resistance, non-pyrophoricity, and mechanical strength.

Compounds

The great majority (about 90%) of aluminium oxide is converted to metallic aluminium.[117] Being a very hard material (Mohs hardness 9),[128] alumina is widely used as an abrasive;[129] being extraordinarily chemically inert, it is useful in highly reactive environments such as high pressure sodium lamps.[130] Aluminium oxide is commonly used as a catalyst for industrial processes;[117] e.g. the Claus process to convert hydrogen sulfide to sulfur in refineries and to alkylate amines.[131][132] Many industrial catalysts are supported by alumina, meaning that the expensive catalyst material is dispersed over a surface of the inert alumina.[133] Another principal use is as a drying agent or absorbent.[117][134]

Laser deposition of alumina on a substrate

Several sulfates of aluminium have industrial and commercial application. Aluminium sulfate (in its hydrate form) is produced on the annual scale of several millions of metric tons.[135] About two-thirds is consumed in water treatment.[135] The next major application is in the manufacture of paper.[135] It is also used as a mordant in dyeing, in pickling seeds, deodorizing of mineral oils, in leather tanning, and in production of other aluminium compounds.[135] Two kinds of alum, ammonium alum and potassium alum, were formerly used as mordants and in leather tanning, but their use has significantly declined following availability of high-purity aluminium sulfate.[135] Anhydrous aluminium chloride is used as a catalyst in chemical and petrochemical industries, the dyeing industry, and in synthesis of various inorganic and organic compounds.[135] Aluminium hydroxychlorides are used in purifying water, in the paper industry, and as antiperspirants.[135] Sodium aluminate is used in treating water and as an accelerator of solidification of cement.[135]

Many aluminium compounds have niche applications, for example:

• Aluminium acetate in solution is used as an astringent.[136]
• Aluminium phosphate is used in the manufacture of glass, ceramic, pulp and paper products, cosmetics, paints, varnishes, and in dental cement.[137]
• Aluminium hydroxide is used as an antacid, and mordant; it is used also in water purification, the manufacture of glass and ceramics, and in the waterproofing of fabrics.[138][139]
• Lithium aluminium hydride is a powerful reducing agent used in organic chemistry.[140][141]
• Organoaluminiums are used as Lewis acids and cocatalysts.[142]
• Methylaluminoxane is a cocatalyst for Ziegler–Natta olefin polymerization to produce vinyl polymers such as polyethene.[143]
• Aqueous aluminium ions (such as aqueous aluminium sulfate) are used to treat against fish parasites such as Gyrodactylus salaris.[144]
• In many vaccines, certain aluminium salts serve as an immune adjuvant (immune response booster) to allow the protein in the vaccine to achieve sufficient potency as an immune stimulant.[145]

Biology

Schematic of aluminium absorption by human skin.[146]

Despite its widespread occurrence in the Earth's crust, aluminium has no known function in biology.[10] At pH 6–9 (relevant for most natural waters), aluminium precipitates out of water as the hydroxide and is hence not available; most elements behaving this way have no biological role or are toxic.[147] Aluminium salts are remarkably nontoxic, aluminium sulfate having an LD₅₀ of 6207 mg/kg (oral, mouse), which corresponds to 435 grams for an 70 kg (150 lb) person.[10]

Toxicity

In most people, aluminium is not as toxic as heavy metals. Aluminium is classified as a non-carcinogen by the United States Department of Health and Human Services.[148] There is little evidence that normal exposure to aluminium presents a risk to healthy adult,[149] and there is evidence of no toxicity if it is consumed in amounts not greater than 40 mg/day per kg of body mass.[148] Most aluminium consumed will leave the body in feces; most of the small part of it that enters the bloodstream, will be excreted via urine.[150]

Effects

Aluminium, although rarely, can cause vitamin D-resistant osteomalacia, erythropoietin-resistant microcytic anemia, and central nervous system alterations. People with kidney insufficiency are especially at a risk.[148] Chronic ingestion of hydrated aluminium silicates (for excess gastric acidity control) may result in aluminium binding to intestinal contents and increased elimination of other metals, such as iron or zinc; sufficiently high doses (>50 g/day) can cause anemia.[148]

There are five major aluminium forms absorbed by human body: the free solvated trivalent cation (Al³⁺_(aq)); low-molecular-weight, neutral, soluble complexes (LMW-Al⁰_(aq)); high-molecular-weight, neutral, soluble complexes (HMW-Al⁰_(aq)); low-molecular-weight, charged, soluble complexes (LMW-Al(L)_n^+/−_(aq)); nano and micro-particulates (Al(L)_n(s)). They are transported across cell membranes or cell epi-/endothelia through five major routes: (1) paracellular; (2) transcellular; (3) active transport; (4) channels; (5) adsorptive or receptor-mediated endocytosis.[146]

During the 1988 Camelford water pollution incident people in Camelford had their drinking water contaminated with aluminium sulfate for several weeks. A final report into the incident in 2013 concluded it was unlikely that this had caused long-term health problems.[151]

Aluminium has been suspected of being a possible cause of Alzheimer's disease,[152] but research into this for over 40 years has found, as of 2018, no good evidence of causal effect.[153][154]

Aluminium increases estrogen-related gene expression in human breast cancer cells cultured in the laboratory.[155] In very high doses, aluminium is associated with altered function of the blood–brain barrier.[156] A small percentage of people[157] have contact allergies to aluminium and experience itchy red rashes, headache, muscle pain, joint pain, poor memory, insomnia, depression, asthma, irritable bowel syndrome, or other symptoms upon contact with products containing aluminium.[158]

Exposure to powdered aluminium or aluminium welding fumes can cause pulmonary fibrosis.[159] Fine aluminium powder can ignite or explode, posing another workplace hazard.[160][161]

Exposure routes

Food is the main source of aluminium. Drinking water contains more aluminium than solid food;[148] however, aluminium in food may be absorbed more than aluminium from water.[162] Major sources of human oral exposure to aluminium include food (due to its use in food additives, food and beverage packaging, and cooking utensils), drinking water (due to its use in municipal water treatment), and aluminium-containing medications (particularly antacid/antiulcer and buffered aspirin formulations).[163] Dietary exposure in Europeans averages to 0.2–1.5 mg/kg/week but can be as high as 2.3 mg/kg/week.[148] Higher exposure levels of aluminium are mostly limited to miners, aluminium production workers, and dialysis patients.[164]

Consumption of antacids, antiperspirants, vaccines, and cosmetics provide possible routes of exposure.[165] Consumption of acidic foods or liquids with aluminium enhances aluminium absorption,[166] and maltol has been shown to increase the accumulation of aluminium in nerve and bone tissues.[167]

Treatment

In case of suspected sudden intake of a large amount of aluminium, the only treatment is deferoxamine mesylate which may be given to help eliminate aluminium from the body by chelation.[168][169] However, this should be applied with caution as this reduces not only aluminium body levels, but also those of other metals such as copper or iron.[168]

Environmental effects

"Bauxite tailings" storage facility in Stade, Germany. The aluminium industry generates about 70 million tons of this waste annually.

High levels of aluminium occur near mining sites; small amounts of aluminium are released to the environment at the coal-fired power plants or incinerators.[170] Aluminium in the air is washed out by the rain or normally settles down but small particles of aluminium remain in the air for a long time.[170]

Acidic precipitation is the main natural factor to mobilize aluminium from natural sources[148] and the main reason for the environmental effects of aluminium;[171] however, the main factor of presence of aluminium in salt and freshwater are the industrial processes that also release aluminium into air.[148]

In water, aluminium acts as a toxiс agent on gill-breathing animals such as fish by causing loss of plasma- and hemolymph ions leading to osmoregulatory failure.[171] Organic complexes of aluminium may be easily absorbed and interfere with metabolism in mammals and birds, even though this rarely happens in practice.[171]

Aluminium is primary among the factors that reduce plant growth on acidic soils. Although it is generally harmless to plant growth in pH-neutral soils, in acid soils the concentration of toxic Al³⁺ cations increases and disturbs root growth and function.[172][173][174][175] Wheat has developed a tolerance to aluminium, releasing organic compounds that bind to harmful aluminium cations. Sorghum is believed to have the same tolerance mechanism.[176]

Aluminium production possesses its own challenges to the environment on each step of the production process. The major challenge is the greenhouse gas emissions.[164] These gases result from electrical consumption of the smelters and the byproducts of processing. The most potent of these gases are perfluorocarbons from the smelting process.[164] Released sulfur dioxide is one of the primary precursors of acid rain.[164]

A Spanish scientific report from 2001 claimed that the fungus Geotrichum candidum consumes the aluminium in compact discs.[177][178] Other reports all refer back to that report and there is no supporting original research. Better documented, the bacterium Pseudomonas aeruginosa and the fungus Cladosporium resinae are commonly detected in aircraft fuel tanks that use kerosene-based fuels (not avgas), and laboratory cultures can degrade aluminium.[179] However, these life forms do not directly attack or consume the aluminium; rather, the metal is corroded by microbe waste products.[180]

See also

• Aluminium granules
• Aluminum joining
• Aluminium–air battery
• Panel edge staining
• Quantum clock

Notes

1. As aluminium technically does not come after any transition metals in the periodic table, it is excluded by some authors from the set of post-transition metals.[3] Nevertheless its weakly metallic behaviour is similar to that of its heavier congeners in group 13 gallium, indium, and thallium, which are post-transition metals by all definitions.
2. No elements with odd atomic numbers have more than two stable isotopes; even-numbered elements have multiple stable isotopes, with tin (element 50) having the highest number of isotopes of all elements, ten.[11] See Even and odd atomic nuclei for more details.
3. Most other metals have greater standard atomic weights: for instance, that of iron is 55.8; copper 63.5; lead 207.2.[1]
4. However, they should not be considered as [AlF₆]³⁻ complex anions as the Al–F bonds are not significantly different in type from the other M–F bonds,[35] and such differences in coordination between the fluorides and heavier halides are not unusual, occurring in Sn^IV and Bi^III, for example; even bigger differences occur between CO₂ and SiO₂.[35]
5. Abundances in the source are listed relative to silicon rather than in per-particle notation. The sum of all elements per 10⁶ parts of silicon is 2.6682×10¹⁰ parts; aluminium comprises 8.410×10⁴ parts.
6. For instance, see the November–December 2013 issue of Chemistry International: in a table of (some) elements, the element is listed as "aluminium (aluminum)".[112]

References

1. Meija, Juris; et al. (2016). "Atomic weights of the elements 2013 (IUPAC Technical Report)". Pure and Applied Chemistry. 88 (3): 265–91. doi:10.1515/pac-2015-0305.
2. Whitten KW, Davis RE, Peck LM & Stanley GG 2014, Chemistry, 10th ed., Thomson Brooks/Cole, Belmont, California, ISBN 1-133-61066-8, p. 1045
3. Cox PA 2004, Inorganic chemistry, 2nd ed., Instant notes series, Bios Scientific, London, ISBN 1-85996-289-0, p. 186
4. Dohmeier, C.; Loos, D.; Schnöckel, H. (1996). "Aluminum(I) and Gallium(I) Compounds: Syntheses, Structures, and Reactions". Angewandte Chemie International Edition. 35 (2): 129–149. doi:10.1002/anie.199601291.
5. D. C. Tyte (1964). "Red (B2Π–A2σ) Band System of Aluminium Monoxide". Nature. 202 (4930): 383. Bibcode:1964Natur.202..383T. doi:10.1038/202383a0.
6. Lide, D. R. (2000). "Magnetic susceptibility of the elements and inorganic compounds" (PDF). CRC Handbook of Chemistry and Physics (81st ed.). CRC Press. ISBN 0849304814.
7. Shakhashiri, B.Z. (17 March 2008). "Chemical of the Week: Aluminum" (PDF). SciFun.org. University of Wisconsin. Archived from the original (PDF) on 9 May 2012. Retrieved 4 March 2012.
8. Singh, Bikram Jit (2014). RSM: A Key to Optimize Machining: Multi-Response Optimization of CNC Turning with Al-7020 Alloy. Anchor Academic Publishing (aap_verlag). ISBN 978-3-95489-209-9.
9. Hihara, Lloyd H.; Adler, Ralph P.I.; Latanision, Ronald M. (2013). Environmental Degradation of Advanced and Traditional Engineering Materials. CRC Press. ISBN 978-1-4398-1927-2.
10. Frank, W.B. (2009). "Aluminum". Ullmann's Encyclopedia of Industrial Chemistry. Wiley-VCH. doi:10.1002/14356007.a01_459.pub2. ISBN 978-3-527-30673-2.
11. IAEA – Nuclear Data Section (2017). "Livechart – Table of Nuclides – Nuclear structure and decay data". www-nds.iaea.org. International Atomic Energy Agency. Retrieved 31 March 2017.
12. Greenwood and Earnshaw, pp. 242–52
13. Dickin, A.P. (2005). "In situ Cosmogenic Isotopes". Radiogenic Isotope Geology. Cambridge University Press. ISBN 978-0-521-53017-0. Archived from the original on 6 December 2008. Retrieved 16 July 2008.
14. Dodd, R.T. (1986). Thunderstones and Shooting Stars. Harvard University Press. pp. 89–90. ISBN 978-0-674-89137-1.
15. Dean 1999, p. 4.2.
16. Dean 1999, p. 4.6.
17. Dean 1999, p. 4.29.
18. Dean 1999, p. 4.30.
19. Enghag, Per (2008). Encyclopedia of the Elements: Technical Data – History – Processing – Applications. John Wiley & Sons. pp. 139, 819, 949. ISBN 978-3-527-61234-5.
20. Greenwood and Earnshaw, pp. 222–4
21. Lide 2004, p. 4-3.
22. Puchta, Ralph (2011). "A brighter beryllium". Nature Chemistry. 3 (5): 416. Bibcode:2011NatCh...3..416P. doi:10.1038/nchem.1033. PMID 21505503.
23. Polmear, I.J. (1995). Light Alloys: Metallurgy of the Light Metals (3 ed.). Butterworth-Heinemann. ISBN 978-0-340-63207-9.
24. Cardarelli, François (2008). Materials handbook : a concise desktop reference (2nd ed.). London: Springer. pp. 158–163. ISBN 978-1-84628-669-8. OCLC 261324602.
25. Cochran, J.F.; Mapother, D.E. (1958). "Superconducting Transition in Aluminum". Physical Review. 111 (1): 132–142. Bibcode:1958PhRv..111..132C. doi:10.1103/PhysRev.111.132.
26. Greenwood and Earnshaw, pp. 224–7
27. Greenwood and Earnshaw, pp. 112–3
28. King, p. 241
29. King, pp. 235–6
30. Hatch, John E. (1984). Aluminum : properties and physical metallurgy. Aluminum Association., American Society for Metals. Metals Park, Ohio: American Society for Metals. p. 242. ISBN 978-1-61503-169-6. OCLC 759213422.
31. Vargel, Christian (2004) [French edition published 1999]. Corrosion of Aluminium. Elsevier. ISBN 978-0-08-044495-6. Archived from the original on 21 May 2016.
32. Macleod, H.A. (2001). Thin-film optical filters. CRC Press. p. 158159. ISBN 978-0-7503-0688-1.
33. Beal, Roy E. (1999). Engine Coolant Testing : Fourth Volume. ASTM International. p. 90. ISBN 978-0-8031-2610-7. Archived from the original on 24 April 2016.
34. • Baes, C.F.; Mesmer, R.E. (1986) [1976]. The Hydrolysis of Cations. Malabar, FL: Robert E. Krieger. ISBN 978-0-89874-892-5.
35. Greenwood and Earnshaw, pp. 233–7
36. Eastaugh, Nicholas; Walsh, Valentine; Chaplin, Tracey; Siddall, Ruth (2008). Pigment Compendium. Routledge. ISBN 978-1-136-37393-0.
37. Roscoe, Henry Enfield; Schorlemmer, Carl (1913). A treatise on chemistry. Macmillan.
38. Greenwood and Earnshaw, pp. 252–7
39. Dohmeier, C.; Loos, D.; Schnöckel, H. (1996). "Aluminum(I) and Gallium(I) Compounds: Syntheses, Structures, and Reactions". Angewandte Chemie International Edition. 35 (2): 129–149. doi:10.1002/anie.199601291.
40. Tyte, D.C. (1964). "Red (B2Π–A2σ) Band System of Aluminium Monoxide". Nature. 202 (4930): 383–384. Bibcode:1964Natur.202..383T. doi:10.1038/202383a0.
41. Merrill, P.W.; Deutsch, A.J.; Keenan, P.C. (1962). "Absorption Spectra of M-Type Mira Variables". The Astrophysical Journal. 136: 21. Bibcode:1962ApJ...136...21M. doi:10.1086/147348.
42. Uhl, W. (2004). "Organoelement Compounds Possessing Al–Al, Ga–Ga, In–In, and Tl–Tl Single Bonds". Organoelement Compounds Possessing Al–Al, Ga–Ga, In–In, and Tl–Tl Single Bonds. Advances in Organometallic Chemistry. 51. pp. 53–108. doi:10.1016/S0065-3055(03)51002-4. ISBN 978-0-12-031151-4.
43. Elschenbroich, C. (2006). Organometallics. Wiley-VCH. ISBN 978-3-527-29390-2.
44. Greenwood and Earnshaw, pp. 257–67
45. Smith, Martin B. (1970). "The monomer-dimer equilibria of liquid aluminum alkyls". Journal of Organometallic Chemistry. 22: 273–281. doi:10.1016/S0022-328X(00)86043-X.
46. Greenwood and Earnshaw, pp. 227–32
47. Lodders, K. (2003). "Solar System abundances and condensation temperatures of the elements" (PDF). The Astrophysical Journal. 591 (2): 1220–1247. Bibcode:2003ApJ...591.1220L. doi:10.1086/375492. ISSN 0004-637X.
48. Clayton, Donald (2007). Handbook of Isotopes in the Cosmos: Hydrogen to Gallium. Cambridge University Press. pp. 129–137. ISBN 978-0-521-53083-5.
49. William F McDonough The composition of the Earth. quake.mit.edu, archived by the Internet Archive Wayback Machine.
50. Greenwood and Earnshaw, pp. 217–9
51. Wade, K.; Banister, A.J. (2016). The Chemistry of Aluminium, Gallium, Indium and Thallium: Comprehensive Inorganic Chemistry. Elsevier. p. 1049. ISBN 978-1-4831-5322-3.
52. Palme, H.; O'Neill, Hugh St. C. (2005). "Cosmochemical Estimates of Mantle Composition" (PDF). In Carlson, Richard W. (ed.). The Mantle and Core. Elseiver. p. 14.
53. Downs, A.J. (1993). Chemistry of Aluminium, Gallium, Indium and Thallium. Springer Science & Business Media. ISBN 978-0-7514-0103-5.
54. Kotz, John C.; Treichel, Paul M.; Townsend, John (2012). Chemistry and Chemical Reactivity. Cengage Learning. p. 300. ISBN 978-1-133-42007-1.
55. Barthelmy, D. "Aluminum Mineral Data". Mineralogy Database. Archived from the original on 4 July 2008. Retrieved 9 July 2008.
56. Chen, Z.; Huang, Chi-Yue; Zhao, Meixun; Yan, Wen; Chien, Chih-Wei; Chen, Muhong; Yang, Huaping; Machiyama, Hideaki; Lin, Saulwood (2011). "Characteristics and possible origin of native aluminum in cold seep sediments from the northeastern South China Sea". Journal of Asian Earth Sciences. 40 (1): 363–370. Bibcode:2011JAESc..40..363C. doi:10.1016/j.jseaes.2010.06.006.
57. Guilbert, J.F.; Park, C.F. (1986). The Geology of Ore Deposits. W.H. Freeman. pp. 774–795. ISBN 978-0-7167-1456-9.
58. United States Geological Survey (2018). "Bauxite and alumina" (PDF). Mineral Commodities Summaries. Retrieved 17 June 2018.
59. Drozdov 2007, p. 12.
60. Clapham, John Harold; Power, Eileen Edna (1941). The Cambridge Economic History of Europe: From the Decline of the Roman Empire. CUP Archive. p. 207. ISBN 978-0-521-08710-0.
61. Drozdov 2007, p. 16.
62. Setton, Kenneth M. (1976). The papacy and the Levant: 1204-1571. 1 The thirteenth and fourteenth centuries. American Philosophical Society. ISBN 978-0-87169-127-9. OCLC 165383496.
63. Drozdov 2007, p. 25.
64. Weeks, Mary Elvira (1968). Discovery of the elements. 1 (7 ed.). Journal of chemical education. p. 187.
65. Richards 1896, p. 2.
66. Richards 1896, p. 3.
67. Örsted, H. C. (1825). Oversigt over det Kongelige Danske Videnskabernes Selskabs Forhanlingar og dets Medlemmerz Arbeider, fra 31 Mai 1824 til 31 Mai 1825 [Overview of the Royal Danish Science Society's Proceedings and the Work of its Members, from 31 May 1824 to 31 May 1825] (in Danish). pp. 15–16.
68. Royal Danish Academy of Sciences and Letters (1827). Det Kongelige Danske Videnskabernes Selskabs philosophiske og historiske afhandlinger [The philosophical and historical dissertations of the Royal Danish Science Society] (in Danish). Popp. pp. xxv–xxvi.
69. Wöhler, Friedrich (1827). "Ueber das Aluminium". Annalen der Physik und Chemie. 2. 11 (9): 146–161. Bibcode:1828AnP....87..146W. doi:10.1002/andp.18270870912.
70. Drozdov 2007, p. 36.
71. Fontani, Marco; Costa, Mariagrazia; Orna, Mary Virginia (2014). The Lost Elements: The Periodic Table's Shadow Side. Oxford University Press. p. 30. ISBN 978-0-19-938334-4.
72. Venetski, S. (1969). "'Silver' from clay". Metallurgist. 13 (7): 451–453. doi:10.1007/BF00741130.
73. Drozdov 2007, p. 38.
74. Holmes, Harry N. (1936). "Fifty Years of Industrial Aluminum". The Scientific Monthly. 42 (3): 236–239. Bibcode:1936SciMo..42..236H. JSTOR 15938.
75. Drozdov 2007, p. 39.
76. Sainte-Claire Deville, H.E. (1859). De l'aluminium, ses propriétés, sa fabrication. Paris: Mallet-Bachelier. Archived from the original on 30 April 2016.
77. Drozdov 2007, p. 46.
78. Drozdov 2007, pp. 55–61.
79. Drozdov 2007, p. 74.
80. "Aluminium history". All about aluminium. Retrieved 7 November 2017.
81. Drozdov 2007, pp. 64–69.
82. Ingulstad, Mats (2012). "'We Want Aluminum, No Excuses': Business-Government Relations in the American Aluminum Industry, 1917–1957". In Ingulstad, Mats; Frøland, Hans Otto (eds.). From Warfare to Welfare: Business-Government Relations in the Aluminium Industry. Tapir Academic Press. pp. 33–68. ISBN 978-82-321-0049-1.
83. Drozdov 2007, pp. 69–70.
84. Drozdov 2007, pp. 165–166.
85. Drozdov 2007, p. 85.
86. Drozdov 2007, p. 135.
87. "Aluminum". Historical Statistics for Mineral Commodities in the United States (Report). United States Geological Survey. 2017. Retrieved 9 November 2017.
88. Nappi 2013, p. 9.
89. Nappi 2013, pp. 9–10.
90. Nappi 2013, p. 10.
91. Nappi 2013, pp. 14–15.
92. Nappi 2013, p. 17.
93. Nappi 2013, p. 20.
94. Nappi 2013, p. 22.
95. Nappi 2013, p. 23.
96. Harper, Douglas. "Alum". Online Etymology Dictionary. Retrieved 13 November 2017.
97. Pokorny, Julius (1959). "alu- (-d-, -t-)". Indogermanisches etymologisches Wörterbuch [Indo-European etymological dictionary] (in German). A. Francke Verlag. pp. 33–34.
98. Davy, Humphry (1808). "Electro Chemical Researches, on the Decomposition of the Earths; with Observations on the Metals obtained from the alkaline Earths, and on the Amalgam procured from Ammonia". Philosophical Transactions of the Royal Society. 98: 353. Bibcode:1808RSPT...98..333D. doi:10.1098/rstl.1808.0023. Retrieved 10 December 2009.
99. Richards 1896, pp. 3–4.
100. "Philosophical Transactions of the Royal Society of London. For the Year 1810. — Part I". The Critical Review: or, Annals of Literature. The Third. XXII: 9. January 1811. Retrieved 9 August 2020 – via Hathi Trust.

     Potassium, acting upon alumine and glucine, produces pyrophoric substances of a dark grey colour, which burnt, throwing off brilliant sparks, and leaving behind alkali and earth, and which, when thrown into water, decomposed it with great violence. The result of this experiment is not wholly decisive as to the existence of what might be called aluminium and glucinium

101. Davy, Humphry (1812). "Of metals; their primary compositions with other uncompounded bodies, and with each other". Elements of Chemical Philosophy: Part 1. 1. Bradford and Inskeep. p. 201.
102. Powell, Mike (2015). Amglish: Two Nations Divided by a Common Language. [Pennsauken, New Jersey]: BookBaby. p. 138. ISBN 978-1-63192-720-1. OCLC 913137419.
103. "-ium, suffix". Oxford English Dictionary (3rd ed.). Oxford University Press. September 2005. Retrieved 8 August 2020. (Subscription or UK public library membership required.)
104. Cutmore, Jonathan (February 2005). "Quarterly Review Archive". Romantic Circles. University of Maryland. Archived from the original on 1 March 2017. Retrieved 28 February 2017.
105. Young, Thomas (1812). Elements of Chemical Philosophy By Sir Humphry Davy. Quarterly Review. VIII. p. 72. ISBN 978-0-217-88947-6. 210. Retrieved 10 December 2009.
106. Quinion, Michael (2005). Port Out, Starboard Home: The Fascinating Stories We Tell About the words We Use. Penguin Books Limited. pp. 23–24. ISBN 978-0-14-190904-2.
107. Richards 1896, p. 4.
108. Webster, Noah (1828). "aluminum". American Dictionary of the English Language. Retrieved 13 November 2017.
109. "Aluminum vs Aluminium". Spectra Aluminum Products. Retrieved 13 November 2017.
110. Emsley, John (2011). Nature's Building Blocks: An A–Z Guide to the Elements. OUP Oxford. pp. 24–30. ISBN 978-0-19-960563-7.
111. Connelly, Neil G.; Damhus, Ture, eds. (2005). Nomenclature of inorganic chemistry. IUPAC Recommendations 2005 (PDF). RSC Publishing. p. 249. ISBN 978-0-85404-438-2. Archived from the original (PDF) on 22 December 2014.
112. "Standard Atomic Weights Revised" (PDF). Chemistry International. 35 (6): 17–18. ISSN 0193-6484. Archived from the original (PDF) on 11 February 2014.
113. Brown, T.J.; Idoine, N.E.; Raycraft, E.R.; et al. (2018). World Mineral Production: 2012–2016. British Geological Survey. ISBN 978-0-85272-882-6.
114. Brown, T.J. (2009). World Mineral Production 2003–2007. British Geological Survey.
115. "Top 10 Largest Aluminium Smelters in the World". Gulf Business. 2013. Retrieved 25 June 2018.
116. Graedel, T.E.; et al. (2010). Metal stocks in Society – Scientific Synthesis (PDF) (Report). International Resource Panel. p. 17. ISBN 978-92-807-3082-1. Retrieved 18 April 2017.
117. Hudson, L. Keith; Misra, Chanakya; Perrotta, Anthony J.; et al. (2005). "Aluminum Oxide". Ullmann's Encyclopedia of Industrial Chemistry. Wiley-VCH.
118. Totten, G.E.; Mackenzie, D.S. (2003). Handbook of Aluminum. Marcel Dekker. p. 40. ISBN 978-0-8247-4843-2. Archived from the original on 15 June 2016.
119. Schlesinger, Mark (2006). Aluminum Recycling. CRC Press. p. 248. ISBN 978-0-8493-9662-5.
120. "Benefits of Recycling". Ohio Department of Natural Resources. Archived from the original on 24 June 2003.
121. "Theoretical/Best Practice Energy Use in Metalcasting Operations" (PDF). Archived from the original (PDF) on 31 October 2013. Retrieved 28 October 2013.
122. "Why are dross & saltcake a concern?". www.experts123.com. Archived from the original on 14 November 2012.
123. Dunster, A.M.; et al. (2005). "Added value of using new industrial waste streams as secondary aggregates in both concrete and asphalt" (PDF). Waste & Resources Action Programme. Archived from the original on 2 April 2010.
124. "Aluminum". Encyclopædia Britannica. Archived from the original on 12 March 2012. Retrieved 6 March 2012.
125. Millberg, L.S. "Aluminum Foil". How Products are Made. Archived from the original on 13 July 2007. Retrieved 11 August 2007.
126. Lyle, J.P.; Granger, D.A.; Sanders, R.E. (2005). "Aluminum Alloys". Ullmann's Encyclopedia of Industrial Chemistry. Wiley-VCH. doi:10.1002/14356007.a01_481. ISBN 978-3-527-30673-2.
127. Davis, Joseph R. (1993). Aluminum and Aluminum Alloys. ASM International. pp. 13–17. ISBN 978-0-87170-496-2.
128. Lumley, Roger (2010). Fundamentals of Aluminium Metallurgy: Production, Processing and Applications. Elsevier Science. p. 42. ISBN 978-0-85709-025-6.
129. Mortensen, Andreas (2006). Concise Encyclopedia of Composite Materials. Elsevier. p. 281. ISBN 978-0-08-052462-7.
130. The Ceramic Society of Japan (2012). Advanced Ceramic Technologies & Products. Springer Science & Business Media. p. 541. ISBN 978-4-431-54108-0.
131. Slesser, Malcolm (1988). Dictionary of Energy. Palgrave Macmillan UK. p. 138. ISBN 978-1-349-19476-6.
132. Supp, Emil (2013). How to Produce Methanol from Coal. Springer Science & Business Media. pp. 164–165. ISBN 978-3-662-00895-9.
133. Ertl, Gerhard; Knözinger, Helmut; Weitkamp, Jens (2008). Preparation of Solid Catalysts. John Wiley & Sons. p. 80. ISBN 978-3-527-62068-5.
134. Armarego, W.L.F.; Chai, Christina (2009). Purification of Laboratory Chemicals. Butterworth-Heinemann. pp. 73, 109, 116, 155. ISBN 978-0-08-087824-9.
135. Helmboldt, O. (2007). "Aluminum Compounds, Inorganic". Ullmann's Encyclopedia of Industrial Chemistry. Wiley-VCH. doi:10.1002/14356007.a01_527.pub2. ISBN 978-3-527-30673-2.
136. World Health Organization (2009). Stuart MC, Kouimtzi M, Hill SR (eds.). WHO Model Formulary 2008. World Health Organization. hdl:10665/44053. ISBN 9789241547659.
137. Occupational Skin Disease. Grune & Stratton. 1983. ISBN 978-0-8089-1494-5.
138. Galbraith, A; Bullock, S; Manias, E; Hunt, B; Richards, A (1999). Fundamentals of pharmacology: a text for nurses and health professionals. Harlow: Pearson. p. 482.
139. Papich, Mark G. (2007). "Aluminum Hydroxide and Aluminum Carbonate". Saunders Handbook of Veterinary Drugs (2nd ed.). St. Louis, Mo: Saunders/Elsevier. pp. 15–16. ISBN 978-1-4160-2888-8.
140. Brown, H.C. (1951). Reductions by Lithium Aluminum Hydride. Organic Reactions. 6. p. 469. doi:10.1002/0471264180.or006.10. ISBN 978-0-471-26418-7.
141. Gerrans, G.C.; Hartmann-Petersen, P. (2007). "Lithium Aluminium Hydride". Sasol Encyclopaedia of Science and Technology. New Africa Books. p. 143. ISBN 978-1-86928-384-1.
142. M. Witt; H.W. Roesky (2000). "Organoaluminum chemistry at the forefront of research and development" (PDF). Curr. Sci. 78 (4): 410. Archived from the original (PDF) on 6 October 2014.
143. A. Andresen; H.G. Cordes; J. Herwig; W. Kaminsky; A. Merck; R. Mottweiler; J. Pein; H. Sinn; H.J. Vollmer (1976). "Halogen-free Soluble Ziegler-Catalysts for the Polymerization of Ethylene". Angew. Chem. Int. Ed. 15 (10): 630–632. doi:10.1002/anie.197606301.
144. Aas, Øystein; Klemetsen, Anders; Einum, Sigurd; et al. (2011). Atlantic Salmon Ecology. John Wiley & Sons. p. 240. ISBN 978-1-4443-4819-4.
145. Singh, Manmohan (2007). Vaccine Adjuvants and Delivery Systems. John Wiley & Sons. pp. 81–109. ISBN 978-0-470-13492-4.
146. Exley, C. (2013). "Human exposure to aluminium". Environmental Science: Processes & Impacts. 15 (10): 1807–1816. doi:10.1039/C3EM00374D. PMID 23982047.
147. "Environmental Applications. Part I. Common Forms of the Elements in Water". Western Oregon University. Western Oregon University. Retrieved 30 September 2019.
148. Dolara, Piero (21 July 2014). "Occurrence, exposure, effects, recommended intake and possible dietary use of selected trace compounds (aluminium, bismuth, cobalt, gold, lithium, nickel, silver)". International Journal of Food Sciences and Nutrition. 65 (8): 911–924. doi:10.3109/09637486.2014.937801. ISSN 1465-3478. PMID 25045935.
149. Physiology of Aluminum in Man. Aluminum and Health. CRC Press. 1988. p. 90. ISBN 0-8247-8026-4. Archived from the original on 19 May 2016.
150. "ATSDR – Public Health Statement: Aluminum". www.atsdr.cdc.gov. Retrieved 18 July 2018.
151. "Lowermoor Water Pollution incident 'unlikely' to have caused long term health effects" (PDF). Committee on Toxicity of Chemicals in Food, Consumer Products and the Environment. 18 April 2013.
152. https://www.researchgate.net/publication/49682395_Aluminum_and_Alzheimer's_Disease_After_a_Century_of_Controversy_Is_there_a_Plausible_Link
153. "Aluminum and dementia: Is there a link?". Alzheimer Society Canada. 24 August 2018.
154. Santibáñez, Miguel; Bolumar, Francisco; García, Ana M (2007). "Occupational risk factors in Alzheimer's disease: a review assessing the quality of published epidemiological studies". Occupational and Environmental Medicine. 64 (11): 723–732. doi:10.1136/oem.2006.028209. ISSN 1351-0711. PMC 2078415. PMID 17525096.
155. Darbre, P.D. (2006). "Metalloestrogens: an emerging class of inorganic xenoestrogens with potential to add to the oestrogenic burden of the human breast". Journal of Applied Toxicology. 26 (3): 191–197. doi:10.1002/jat.1135. PMID 16489580.
156. Banks, W.A.; Kastin, A.J. (1989). "Aluminum-induced neurotoxicity: alterations in membrane function at the blood–brain barrier". Neurosci Biobehav Rev. 13 (1): 47–53. doi:10.1016/S0149-7634(89)80051-X. PMID 2671833.
157. Bingham, Eula; Cohrssen, Barbara (2012). Patty's Toxicology, 6 Volume Set. John Wiley & Sons. p. 244. ISBN 978-0-470-41081-3.
158. "Aluminum Allergy Symptoms and Diagnosis". Allergy-symptoms.org. 20 September 2016. Retrieved 23 July 2018.
159. al-Masalkhi, A.; Walton, S.P. (1994). "Pulmonary fibrosis and occupational exposure to aluminum". The Journal of the Kentucky Medical Association. 92 (2): 59–61. ISSN 0023-0294. PMID 8163901.
160. "CDC – NIOSH Pocket Guide to Chemical Hazards – Aluminum". www.cdc.gov. Archived from the original on 30 May 2015. Retrieved 11 June 2015.
161. "CDC – NIOSH Pocket Guide to Chemical Hazards – Aluminum (pyro powders and welding fumes, as Al)". www.cdc.gov. Archived from the original on 30 May 2015. Retrieved 11 June 2015.
162. Yokel R.A.; Hicks C.L.; Florence R.L. (2008). "Aluminum bioavailability from basic sodium aluminum phosphate, an approved food additive emulsifying agent, incorporated in cheese". Food and Chemical Toxicology. 46 (6): 2261–2266. doi:10.1016/j.fct.2008.03.004. PMC 2449821. PMID 18436363.
163. United States Department of Health and Human Services (1999). Toxicological profile for aluminum (PDF) (Report). Retrieved 3 August 2018.
164. "Aluminum". The Environmental Literacy Council. Retrieved 29 July 2018.
165. Chen, Jennifer K.; Thyssen, Jacob P. (2018). Metal Allergy: From Dermatitis to Implant and Device Failure. Springer. p. 333. ISBN 978-3-319-58503-1.
166. Slanina, P.; French, W.; Ekström, L.G.; Lööf, L.; Slorach, S.; Cedergren, A. (1986). "Dietary citric acid enhances absorption of aluminum in antacids". Clinical Chemistry. 32 (3): 539–541. doi:10.1093/clinchem/32.3.539. PMID 3948402.
167. Van Ginkel, M.F.; Van Der Voet, G.B.; D'haese, P.C.; De Broe, M.E.; De Wolff, F.A. (1993). "Effect of citric acid and maltol on the accumulation of aluminum in rat brain and bone". The Journal of Laboratory and Clinical Medicine. 121 (3): 453–460. PMID 8445293.
168. "ARL: Aluminum Toxicity". www.arltma.com. Retrieved 24 July 2018.
169. Aluminum Toxicity Archived 3 February 2014 at the Wayback Machine from NYU Langone Medical Center. Last reviewed November 2012 by Igor Puzanov, MD
170. "ATSDR – Public Health Statement: Aluminum". www.atsdr.cdc.gov. Retrieved 28 July 2018.
171. Rosseland, B.O.; Eldhuset, T.D.; Staurnes, M. (1990). "Environmental effects of aluminium". Environmental Geochemistry and Health. 12 (1–2): 17–27. doi:10.1007/BF01734045. ISSN 0269-4042. PMID 24202562.
172. Belmonte Pereira, Luciane; Aimed Tabaldi, Luciane; Fabbrin Gonçalves, Jamile; Jucoski, Gladis Oliveira; Pauletto, Mareni Maria; Nardin Weis, Simone; Texeira Nicoloso, Fernando; Brother, Denise; Batista Teixeira Rocha, João; Chitolina Schetinger, Maria Rosa Chitolina (2006). "Effect of aluminum on δ-aminolevulinic acid dehydratase (ALA-D) and the development of cucumber (Cucumis sativus)". Environmental and Experimental Botany. 57 (1–2): 106–115. doi:10.1016/j.envexpbot.2005.05.004.
173. Andersson, Maud (1988). "Toxicity and tolerance of aluminium in vascular plants". Water, Air, & Soil Pollution. 39 (3–4): 439–462. doi:10.1007/BF00279487 (inactive 29 April 2020).
174. Horst, Walter J. (1995). "The role of the apoplast in aluminium toxicity and resistance of higher plants: A review". Zeitschrift für Pflanzenernährung und Bodenkunde. 158 (5): 419–428. doi:10.1002/jpln.19951580503.
175. Ma, Jian Feng; Ryan, P.R.; Delhaize, E. (2001). "Aluminium tolerance in plants and the complexing role of organic acids". Trends in Plant Science. 6 (6): 273–278. doi:10.1016/S1360-1385(01)01961-6. PMID 11378470.
176. Magalhaes, J.V.; Garvin, D.F.; Wang, Y.; Sorrells, M.E.; Klein, P.E.; Schaffert, R.E.; Li, L.; Kochian, L.V. (2004). "Comparative Mapping of a Major Aluminum Tolerance Gene in Sorghum and Other Species in the Poaceae". Genetics. 167 (4): 1905–1914. doi:10.1534/genetics.103.023580. PMC 1471010. PMID 15342528.
177. "Fungus 'eats' CDs". BBC. 22 June 2001. Archived from the original on 12 December 2013.
178. Bosch, Xavier (27 June 2001). "Fungus eats CD". Nature. doi:10.1038/news010628-11. Archived from the original on 31 December 2010.
179. Sheridan, J.E.; Nelson, Jan; Tan, Y.L. "Studies on the 'Kerosene Fungus' Cladosporium resinae (Lindau) De Vries: Part I. The Problem of Microbial Contamination of Aviation Fuels". Tuatara. 19 (1): 29. Archived from the original on 13 December 2013.
180. "Fuel System Contamination & Starvation". Duncan Aviation. 2011. Archived from the original on 25 February 2015.

Bibliography

• Drozdov, Andrey (2007). Aluminium: The Thirteenth Element. RUSAL Library. ISBN 978-5-91523-002-5.CS1 maint: ref=harv (link)
• Greenwood, Norman N.; Earnshaw, Alan (1997). Chemistry of the Elements (2nd ed.). Butterworth-Heinemann. ISBN 978-0-08-037941-8.
• King, R. Bruce (1995). Inorganic Chemistry of Main Group Elements. Wiley-VCH. ISBN 978-0-471-18602-1.
• Dean, John A. (1999). Lange's handbook of chemistry (15 ed.). McGraw-Hill. ISBN 978-0-07-016384-3. OCLC 40213725.
• Lide, David R., ed. (2004). Handbook of Chemistry and Physics (84 ed.). CRC Press. ISBN 978-0-8493-0566-5.
• Nappi, Carmine (2013). The global aluminium industry 40 years from 1972 (PDF) (Report). International Aluminium Institute. Retrieved 10 November 2017.
• Richards, Joseph William (1896). Aluminium: Its history, occurrence, properties, metallurgy and applications, including its alloys (3 ed.). Henry Carey Baird & Co.

Further reading

• Mimi Sheller, Aluminum Dream: The Making of Light Modernity. Cambridge, Mass.: Massachusetts Institute of Technology Press, 2014.

External links

• Aluminium at The Periodic Table of Videos (University of Nottingham)
• Toxic Substances Portal – Aluminum – from the Agency for Toxic Substances and Disease Registry, United States Department of Health and Human Services
• CDC – NIOSH Pocket Guide to Chemical Hazards – Aluminum
• World production of primary aluminium, by country
• Price history of aluminum, according to the IMF
• History of Aluminium – from the website of the International Aluminium Institute
• Emedicine – Aluminium
• The short film Aluminum(1941) is available for free download at the Internet Archive