Aluminium

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Aluminium, 13Al
Aluminium-4.jpg
Aluminium
Pronunciation
aluminium: /?�lj?'m?ni?m/ (About this soundlisten)
(AL-yuu-MIN-ee-?m)
aluminum: /?'lju?m?n?m/ (About this soundlisten)
(?-LEW-min-?m)
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][a] sometimes considered a metalloid
Electron configuration [Ne] 3s2 3p1
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/cm3
when liquid (at m.p.) 2.375 g/cm3
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)Face-centered cubic crystal structure
for aluminium
Speed of sound thin rod (rolled) 5000 m/s (at r.t.)
Thermal expansion 23.1 �m/(m�K) (at 25 �C)
Thermal conductivity 237 W/(m�K)
Electrical resistivity 26.5 nO�m (at 20 �C)
Magnetic ordering paramagnetic[6]
Magnetic susceptibility +16.5�10-6 cm3/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 and first isolation Hans Christian �rsted (1824)
Named by Humphry Davy (1812)
Main isotopes of aluminium
Iso�tope Abun�dance Half-life (t1/2) Decay mode Pro�duct
26Al trace 7.17�105 y �+ 26Mg
e 26Mg
? �
27Al 100% stable
viewtalkedit | references
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; it is the third most abundant element after oxygen and silicon and
the most abundant metal in the crust, though it is less common in the mantle below.
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]

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.

Contents
1 Physical characteristics
1.1 Nuclei and isotopes
1.2 Electron shell
1.3 Bulk
2 Chemistry
2.1 Inorganic compounds
2.1.1 Rarer oxidation states
2.1.1.1 Aluminium(I)
2.1.1.2 Aluminium(II)
2.2 Organoaluminium compounds and related hydrides
3 Natural occurrence
3.1 In space
3.2 On Earth
4 History
5 Etymology
5.1 Spelling
6 Production and refinement
6.1 Bayer process
6.2 Hall�H�roult process
6.3 Recycling
7 Applications
7.1 Metal
7.2 Compounds
8 Biology
8.1 Toxicity
8.2 Effects
8.3 Exposure routes
8.4 Treatment
9 Environmental effects
10 See also
11 Notes
12 References
13 Bibliography
14 Further reading
15 External links
Physical characteristics
Nuclei and isotopes
Main article: Isotopes of aluminium
Of aluminium isotopes, only 27
Al
is stable. This is consistent with aluminium having an odd atomic number.[b] It is
the only aluminium isotope that has existed on Earth in its current form since the
creation of the planet. Very nearly all the element on Earth is present as this
isotope, which makes aluminium a mononuclidic element and means that its standard
atomic weight practically equates to that of the isotope. The standard atomic
weight of aluminium is low in comparison with many other metals,[c] 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
(half-life 720,000 years) and therefore 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]3s23p1,[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]

M. Tunes & S. Pogatscher, Montanuniversit�t Leoben 2019 No copyrights =)

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 face-centered cubic 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 is 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, 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).[24] It is paramagnetic and thus essentially unaffected by static
magnetic fields. The high electrical conductivity, however, means that it is
strongly affected by changing magnetic field through the induction of eddy
currents.

Etched surface from a high purity (99.9998%) aluminium bar, size 55�37 mm
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;[25] this behaviour is similar to that of beryllium (Be2+), an
example of a diagonal relationship.[26] However, unlike all other post-transition
metals, the underlying core under aluminium's valence shell is that of the
preceding noble gas, whereas for gallium and indium it is that of the preceding
noble gas plus a filled d-subshell, and for thallium and nihonium it is that of the
preceding noble gas plus filled d- and f-subshells. Hence, aluminium does not
suffer the effects of incomplete shielding of valence electrons by inner electrons
from the nucleus that its heavier congeners do. 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 have ds2 configurations of three valence electrons outside a noble
gas core: aluminium is the most electropositive metal in its group.[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 from adducts.[27]
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.[28]

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. 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
that protects the metal from further corrosion by oxygen, water, or dilute acid, a
process termed passivation.[25][29] This layer is destroyed by contact with mercury
due to amalgamation or with salts of some electropositive metals.[25] 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] In addition,
although the reaction of aluminium with water at temperatures below 280 �C is of
interest for the production of hydrogen, commercial application of this fact has
challenges in circumventing the passivating oxide layer, which inhibits the
reaction, and in storing the energy required to regenerate the aluminium metal.[30]

Primarily because it is corroded by dissolved chlorides, such as common sodium

chloride, household plumbing is never made from aluminium.[31] However, 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. Aluminium mirror finish has the highest
reflectance of any metal in the 200�400 nm (UV) and the 3,000�10,000 nm (far IR)
regions; in the 400�700 nm visible range it is slightly outperformed by tin and
silver and in the 700�3000 nm (near IR) by silver, gold, and copper.[32]

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.[31] The reaction with aqueous alkali is often written:[25]

Al + NaOH + H2O ? NaAlO2 +

3
/
2
H2
although the aluminium species in solution is probably instead the hydrated
tetrahydroxoaluminate anion, [Al(OH)4]- or [Al(H2O)2(OH)4]-.[25]

Oxidizing acids do not effectively attack high-purity aluminium because an oxide

layer forms and protects the metal; aqua regia will nevertheless dissolve
aluminium. This allows aluminium to be used to store reagents such as nitric acid,
concentrated sulfuric acid, and some organic acids.[10]

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.[25]

Aluminium hydrolysis as a function of pH. Coordinated water molecules are omitted.

(Data from Baes and Mesmer)[33]
In aqueous solution, Al3+ exists as the hexaaqua cation [Al(H2O)6]3+, which has an
approximate pKa of 10-5.[12] Such solutions are acidic as this cation can act as a
proton donor, progressively hydrolysing to [Al(H2O)5(OH)]2+, [Al(H2O)4(OH)2]+, and
so on. As pH increases these mononuclear species begin to aggregate together by the
formation of hydroxide bridges,[25] forming many oligomeric ions, such as the
Keggin ion [Al13O4(OH)24(H2O)12]7+.[12] The process ends with precipitation of
aluminium hydroxide, Al(OH)3. 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:[25]

Al2O3 + 3 SiO2
fuse
?

Al2(SiO3)3
Al2O3 + CaO
fuse
?

Ca(AlO2)2
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. Further examples include Be2+, Zn2+, Ga3+, Sn2+, and Pb2+;
indeed, gallium in the same group is slightly more acidic than aluminium. One
effect of this is that aluminium salts with weak acids are hydrolysed in water to
the aquated hydroxide and the corresponding nonmetal hydride: aluminium sulfide
yields hydrogen sulfide, aluminium nitride yields ammonia, and aluminium carbide
yields methane. Aluminium cyanide, acetate, and carbonate exist in aqueous solution
but are unstable as such; 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:[25]

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 in a structure related to but distorted from that of ReO3. Such {AlF6}
units also exist in complex fluorides such as cryolite, Na3AlF6, but 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.[34] Such differences in coordination
between the fluorides and heavier halides are not unusual, occurring in SnIV and
BiIII as well for example; even bigger differences occur between CO2 and SiO2.[34]
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).[34]

Mechanism of the Friedel�Crafts acylation, using AlCl3 as a catalyst

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 with a concomitant
increase in volume by 85% and a near-total loss of electrical conductivity. These
still predominate in the gas phase at low temperatures (150�200 �C), but at higher
temperatures 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.[34] 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).[34]

AlCl3 + 3 LiZ ? 3 LiCl + AlZ3 (Z = R, NR2, N=CR2)

AlCl3 + 4 LiZ ? 3 LiCl + LiAlZ4 (Z = R, NR2, N=CR2, H)
BF3 + AlCl3 ? AlF3 + BCl3
Aluminium forms one stable oxide with the chemical formula Al2O3, commonly called
alumina.[35] It can be found in nature in the mineral corundum, a-alumina;[36]
there is also a ?-alumina phase.[12] As 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 cermanics, 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-�-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 (a) or random (�) fashion; the sulfide also has a ?
form related to ?-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.[37] 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.[37]

Rarer oxidation states

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.

Aluminium(I)
Main article: Aluminium(I)
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.[34]

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.[38]

Aluminium(II)
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[39] and in stellar absorption spectra.[40] More thoroughly
investigated are compounds of the formula R4Al2 which contain an Al�Al bond and
where R is a large organic ligand.[41]

Organoaluminium compounds and related hydrides

Main article: Organoaluminium compound

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

A variety of compounds of empirical formula AlR3 and AlR1.5Cl1.5 exist.[42] 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.[43][44] 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.[43]

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:[45]

4 LiH + AlCl3 ? LiAlH4 + 3 LiCl

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.[45]

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).[46][d] 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.[46] 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.[47] 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.[47]

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).
[48] 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.[47] In the
Earth's crust, aluminium is the most abundant (8.3% by mass) metallic element and
the third most abundant of all elements (after oxygen and silicon).[49] A large
number of silicates in the Earth's crust contain aluminium.[50] In contrast, the
Earth's mantle is only 2.38% aluminium by mass.[51]

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.[52]
Impurities in Al2O3, such as chromium and iron, yield the gemstones ruby and
sapphire, respectively.[53] Native aluminium metal can only be found as a minor
phase in low oxygen fugacity environments, such as the interiors of certain
volcanoes.[54] 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-.[55]

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 (AlOx(OH)3�2x). Bauxite occurs as a weathering
product of low iron and silica bedrock in tropical climatic conditions.[56] In
2017, most bauxite was mined in Australia, China, Guinea, and India.[57]

History
Main article: History of aluminium
Friedrich W�hler, 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.[58]
The ancients are known to have used alum as a dyeing mordant and for city defense.
[58] After the Crusades, alum, an indispensable good in the European fabric
industry,[59] was a subject of international commerce;[60] it was imported to
Europe from the eastern Mediterranean until the mid-15th century.[61]

The nature of alum remained unknown. Around 1530, Swiss physician Paracelsus
suggested alum was a salt of an earth of alum.[62] In 1595, German doctor and
chemist Andreas Libavius experimentally confirmed this.[63] In 1722, German chemist
Friedrich Hoffmann announced his belief that the base of alum was a distinct earth.
[64] In 1754, German chemist Andreas Sigismund Marggraf synthesized alumina by
boiling clay in sulfuric acid and subsequently adding potash.[64]

Attempts to produce aluminium metal date back to 1760.[65] The first successful
attempt, however, was completed in 1824 by Danish physicist and chemist Hans
Christian �rsted. He reacted anhydrous aluminium chloride with potassium amalgam,
yielding a lump of metal looking similar to tin.[66][67] He presented his results
and demonstrated a sample of the new metal in 1825.[68][69] In 1827, German chemist
Friedrich W�hler repeated �rsted's experiments but did not identify any aluminium.
[70] (The reason for this inconsistency was only discovered in 1921.)[71] He
conducted a similar experiment in the same year by mixing anhydrous aluminium
chloride with potassium and produced a powder of aluminium.[67] In 1845, he was
able to produce small pieces of the metal and described some physical properties of
this metal.[71] For many years thereafter, W�hler was credited as the discoverer of
aluminium.[72]

The statue of Anteros in Piccadilly Circus, London, was made in 1893 and is one of
the first statues cast in aluminium.
As W�hler's method could not yield great quantities of aluminium, the metal
remained rare; its cost exceeded that of gold.[70] The first industrial production
of aluminium was established in 1856 by French chemist Henri Etienne Sainte-Claire
Deville and companions.[73] Deville had discovered that aluminium trichloride could
be reduced by sodium, which was more convenient and less expensive than potassium,
which W�hler had used.[74] Even then, aluminium was still not of great purity and
produced aluminium differed in properties by sample.[75]

The first industrial large-scale production method was independently developed in

1886 by French engineer Paul H�roult and American engineer Charles Martin Hall; it
is now known as the Hall�H�roult process.[76] The Hall�H�roult 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.[77]
Modern production of the aluminium metal is based on the Bayer and Hall�H�roult
processes.[78]

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.[79] During World War I, major
governments demanded large shipments of aluminium for light strong airframes.[80]

By the mid-20th century, aluminium had become a part of everyday life and an
essential component of housewares.[81] During the mid-20th century, aluminium
emerged as a civil engineering material, with building applications in both basic
construction and interior finish work,[82] and increasingly being used in military
engineering, for both airplanes and land armor vehicle engines.[83] 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.[78] The aluminium can was invented in 1956 and employed as a storage
for drinks in 1958.[84]

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.[85] 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.[78] The output continued to grow: the annual
production of aluminium exceeded 50,000,000 metric tons in 2013.[85]

The real price for aluminium declined from $14,000 per metric ton in 1900 to $2,340
in 1948 (in 1998 United States dollars).[85] 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;[86] the real
price began to grow in the 1970s with the rise of energy cost.[87] Production moved
from the industrialized countries to countries where production was cheaper.[88]
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.[89] The BRIC countries' combined share in primary production and
primary consumption grew substantially in the first decade of the 21st century.[90]
China is accumulating an especially large share of world's production thanks to
abundance of resources, cheap energy, and governmental stimuli;[91] it also
increased its consumption share from 2% in 1972 to 40% in 2010.[92] In the United
States, Western Europe, and Japan, most aluminium was consumed in transportation,
engineering, construction, and packaging.[93]

Etymology
Aluminium is named after alumina, or aluminium oxide in modern nomenclature. The
word "alumina" comes from "alum", the mineral from which it was collected. The word
"alum" comes from alumen, a Latin word meaning "bitter salt".[94] The word alumen
stems from the Proto-Indo-European root *alu- meaning "bitter" or "beer".[95]

1897 American advertisement featuring the aluminum spelling

British chemist Humphry Davy, who performed a number of experiments aimed to
synthesize the metal, is credited as the person who named the element. In 1808, he
suggested the metal be named alumium.[96] 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.[97] In
1812, Davy chose aluminum, thus producing the modern name.[98] However, its
spelling and pronunciation varies: aluminum is in use in the United States and
Canada while aluminium is in use elsewhere.[99]

Spelling
The -ium suffix followed the precedent set in other newly discovered elements of
the time: potassium, sodium, magnesium, calcium, and strontium (all of which Davy
isolated himself). Nevertheless, element names ending in -um were known at the
time; for example, platinum (known to Europeans since the 16th century), molybdenum
(discovered in 1778), and tantalum (discovered in 1802). The -um suffix is
consistent with the universal spelling alumina for the oxide (as opposed to
aluminia); compare to lanthana, the oxide of lanthanum, and magnesia, ceria, and
thoria, the oxides of magnesium, cerium, and thorium, respectively.
In 1812, British scientist Thomas Young[100] wrote an anonymous review of Davy's
book, in which he objected to aluminum and proposed the name aluminium: "for so we
shall take the liberty of writing the word, in preference to aluminum, which has a
less classical sound."[101] This name did catch on: while the -um spelling was
occasionally used in Britain, the American scientific language used -ium from the
start.[102] Most scientists used -ium throughout the world in the 19th century;
[103] it still remains the standard in most other languages.[99] In 1828, American
lexicographer Noah Webster used exclusively the aluminum spelling in his American
Dictionary of the English Language.[104] 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.[102] 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.
[99] 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.[105] In 1925, the American Chemical Society
adopted this spelling.[105]

The International Union of Pure and Applied Chemistry (IUPAC) adopted aluminium as
the standard international name for the element in 1990.[106] In 1993, they
recognized aluminum as an acceptable variant;[106] the most recent 2005 edition of
the IUPAC nomenclature of inorganic chemistry acknowledges this spelling as well.
[107] IUPAC official publications use the -ium spelling as primary but lists both
where appropriate.[e]

Production and refinement