aluminum

—aluminic /al'yeuh min"ik/, adj.

/euh looh"meuh neuhm/, n.

1. Chem. a silver-white metallic element, light in weight, ductile, malleable, and not readily corroded or tarnished, occurring combined in nature in igneous rock, shale, clay, and most soil: used in alloys and for lightweight utensils, castings, airplane parts, etc. Abbr.: alum.; Symbol: Al; at. wt.: 26.98; at. no.: 13; sp. gr.: 2.70 at 20°C.

adj.

2. of, pertaining to, or containing aluminum: an aluminum frying pan. Also, esp. Brit., aluminium.

[1812; < NL, alter., by Humphry Davy, of alumium, which was first proposed; aluminium formed after other metals in -ium. See ALUMINA, -IUM]

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Metallic chemical element, chemical symbol Al, atomic number 13.

A lightweight, silvery white metal, it is so reactive chemically that it always occurs in compounds. It is the most abundant metallic element in Earth's crust, chiefly in bauxite (its principal ore), feldspars, micas, clay minerals, and laterite. It also occurs in gemstones, such as topaz, garnet, and chrysoberyl; emery, corundum, ruby, and sapphire are crystalline aluminum oxide. Aluminum was first isolated in 1825, became commercially available in the late 19th century, and is now the most widely used metal after iron. Its surface oxidizes at once to a hard, tough film, deterring further corrosion. Uses include building and construction, corrosion-resistant chemical equipment, auto and aircraft parts, power transmission lines, photoengraving plates, cookware and other consumer goods, and tubes for ointments and pastes. Important compounds include alums; alumina (aluminum oxide), useful as corundum and as a carrier for many catalysts; aluminum chloride, a widely used catalyst for organic syntheses; and aluminum hydroxide, used to waterproof fabrics.

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▪ chemical element

Introduction

also spelled  aluminium 

 chemical element, a lightweight, silvery-white metal of main Group 13 (IIIa, or boron group) of the periodic table. Aluminum is the most abundant metallic element in the Earth's crust and the most widely used nonferrous metal. Because of its chemical activity, aluminum never occurs in the metallic form in nature, but its compounds are present to a greater or lesser extent in almost all rocks, vegetation, and animals. Aluminum is concentrated in the outer 10 miles (16 km) of the Earth's crust, of which it constitutes about 8 percent by weight; it is exceeded in amount only by oxygen and silicon. The name aluminum is derived from the Latin word alumen, used to describe potash alum, or aluminum potassium sulfate, KAl(SO₄)₂∙12H₂O.

Occurrence, uses, and properties

      Aluminum occurs in igneous rocks chiefly as aluminosilicates in feldspars, feldspathoids, and micas; in the soil derived from them as clay; and upon further weathering as bauxite and iron-rich laterite. Bauxite, a mixture of hydrated aluminum oxides, is the principal aluminum ore. Crystalline aluminum oxide (emery, corundum), which occurs in a few igneous rocks, is mined as a natural abrasive or in its finer varieties as rubies and sapphires. Aluminum is present in other gemstones, such as topaz, garnet, and chrysoberyl. Of the many other aluminum minerals, alunite and cryolite have some commercial importance.

      Crude aluminum was isolated (1825) by Hans Christian Ørsted (Ørsted, Hans Christian) by reducing aluminum chloride with potassium amalgam. Sir Humphry Davy (Davy, Sir Humphry, Baronet) had prepared (1809) an iron-aluminum alloy by electrolyzing fused alumina (aluminum oxide) and had already named the element aluminum; the word later was modified to aluminium in England and some other European countries. A German chemist, Friedrich Wöhler (Wöhler, Friedrich), using potassium metal as the reducing agent, produced aluminum powder (1827) and small globules of the metal (1845) from which he was able to determine some of its properties.

      The new metal was introduced to the public (1855) at the Paris Exposition at about the time that it became available (in small amounts at great expense) by the sodium reduction of molten aluminum chloride. When electric power became relatively plentiful and cheap, almost simultaneously Charles Martin Hall (Hall, Charles Martin) in the United States and Paul-Louis-Toussaint Héroult (Héroult, Paul-Louis-Toussaint) in France discovered (1886) the modern method of commercially producing aluminum: electrolysis of purified alumina (Al₂O₃) dissolved in molten cryolite (Na₃AlF₆). During the 1960s aluminum moved into first place, ahead of copper, in world production of nonferrous metals. For more specific information about the mining, refining, and production of aluminum, see aluminum processing.

      Aluminum is added in small amounts to certain metals to improve their properties for specific uses, as in aluminum bronzes and most magnesium-base alloys; or, for aluminum-base alloys (alloy), moderate amounts of other metals and silicon are added to aluminum. The metal and its alloys are used extensively for aircraft construction, building materials, consumer durables (refrigerators, air conditioners, cooking utensils), electrical conductors, and chemical and food-processing equipment.

      Pure aluminum (99.996 percent) is quite soft and weak; commercial aluminum (99.0 to 99.6 percent pure) with small amounts of silicon and iron is hard and strong. Ductile and highly malleable, aluminum can be drawn into wire or rolled into thin foil. The metal is only about one-third as dense as iron or copper. Though chemically active, aluminum is nevertheless highly corrosion-resistant because in air a hard, tough oxide film forms on its surface.

      Aluminum is an excellent conductor of heat and electricity. Its thermal conductivity is about one-half that of copper; its electrical conductivity, about two-thirds. It crystallizes in the face-centred cubic structure. All natural aluminum is the stable isotope aluminum-27. Metallic aluminum and its oxide and hydroxide are nontoxic.

      Aluminum is slowly attacked by most dilute acids and rapidly dissolves in concentrated hydrochloric acid. Concentrated nitric acid, however, can be shipped in aluminum tank cars because it renders the metal passive. Even very pure aluminum is vigorously attacked by alkalies such as sodium and potassium hydroxide to yield hydrogen and the aluminate ion. Because of its great affinity for oxygen, finely divided aluminum, if ignited, will burn in carbon monoxide or carbon dioxide with the formation of aluminum oxide and carbide; but, at temperatures up to red heat, aluminum is inert to sulfur.

      Aluminum can be detected in concentrations as low as one part per million using emission spectroscopy. Aluminum can be quantitatively analyzed as the oxide (formula Al₂O₃) or as a derivative of the organic nitrogen compound 8-hydroxyquinoline. The derivative has the molecular formula Al(C₉H₆ON)₃.

Compounds

      Ordinarily, aluminum is trivalent. At elevated temperatures, however, a few gaseous monovalent and bivalent compounds have been prepared (AlCl, Al₂O, AlO). In aluminum the configuration of the three outer electrons is such that in a few compounds (e.g., crystalline aluminum fluoride [AlF₃] and aluminum chloride [AlCl₃]) the bare ion, Al³⁺, formed by loss of these electrons, is known to occur. The energy required to form the Al³⁺ ion, however, is very high; and, in the majority of cases, it is energetically more favourable for the aluminum atom to form covalent compounds by way of sp² hybridization, as boron does. The Al³⁺ ion can be stabilized by hydration, and the octahedral ion [Al(H₂O)₆]³⁺ occurs both in aqueous solution and in several salts.

      A number of aluminum compounds have important industrial applications. alumina, which occurs in nature as corundum, is also prepared commercially in large quantities for use in the production of aluminum metal and the manufacture of insulators, spark plugs, and various other products. Upon heating, alumina develops a porous structure, which enables it to adsorb water vapour. This form of aluminum oxide, commercially known as activated alumina, is used for drying gases and certain liquids. It also serves as a carrier for catalysts of various chemical reactions.

      Another major compound is aluminum sulfate, a colourless salt obtained by the action of sulfuric acid on hydrated aluminum oxide. The commercial form is a hydrated crystalline solid with the chemical formula Al₂(SO₄)₃. It is used extensively in paper manufacture as a binder for dyes and as a surface filler. Aluminum sulfate combines with the sulfates of univalent metals to form hydrated double sulfates called alums (alum). The alums, double salts of formula MAl(SO₄)₂ ·12H₂O (where M is a singly charged cation such as K⁺), also contain the Al³⁺ ion; M can be the cation of sodium, potassium, rubidium, cesium, ammonium, or thallium, and the aluminum may be replaced by a variety of other M³⁺ ions: e.g., gallium, indium, titanium, vanadium, chromium, manganese, iron, or cobalt. The most important of such salts is aluminum potassium sulfate, also known as potassium alum or potash alum. These alums have many applications, especially in the production of medicines, textiles, and paints.

      The reaction of gaseous chlorine with molten aluminum metal produces aluminum chloride; the latter is the most commonly used catalyst in Friedel-Crafts reactions—i.e., synthetic organic reactions involved in the preparations of a wide variety of compounds, including aromatic ketones and anthroquinone and its derivatives. Hydrated aluminum chloride, commonly known as aluminum chlorohydrate, AlCl₃∙H₂O, is used as a topical antiperspirant or body deodorant, which acts by constricting the pores. It is one of several aluminum salts employed by the cosmetics industry.

      Aluminum hydroxide, Al(OH)₃, is used to waterproof fabrics and to produce a number of other aluminum compounds, including salts called aluminates that contain the AlO^-₂ group. With hydrogen, aluminum forms aluminum hydride, AlH₃, a polymeric solid from which are derived the tetrohydroaluminates (important reducing agents). Lithium aluminum hydride (LiAlH₄), formed by the reaction of aluminum chloride with lithium hydride, is widely used in organic chemistry—e.g., to reduce aldehydes and ketones to primary and secondary alcohols, respectively.

atomic number

13

atomic weight

26.9815

melting point

660° C

boiling point

2,467° C

specific gravity

2.70 (20° C)

valence

3

electronic config.

2-8-3 or 1s²2s²2p⁶3s²3p¹

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