The symbol in the title is Dalton's symbol for alumina, which was thought to be an element at the time.
This article demonstrates how interesting information can be turned up by researching a chemical element or important compound. It merely presents some entertaining lore, without attempting to be comprehensive. Useful properties of aluminum are listed here in one place, which may be found convenient, and a few topics are followed up in some detail when they demonstrate a general principle. The technical aspects of aluminium as a material of construction, such things as alloys, heat treatment, machining and so forth, are not considered here.
You probably noted that the title uses "aluminium" instead of the American "aluminum," which I did purely in futile protest. Until 1925, the word was "aluminium" even in the U.S., but in that year the American Chemical Society decided to change it. We also got "sulfur" in that same year, which still looks silly, and was not universally adopted by the engineering world. It's the Latin spelling, as is "sulpur." Fortunately, the urge for simplified spelling did not result in Fosforus or Thorum, or even Jermanum, combining both types of change. The -ia ending of a refractory oxide, such as alumina or thoria, usually named the metal with an -ium ending. Why aluminum had to be different, I do not know. A divergence in pronunciation also results, "alyouminium" versus "aloominum." The latter may have been a vulgar pronunciation. It is usually the English who have trouble pronouncing more than three syllables in a word, not the colonials.
Aluminium has the chemical symbol Al, atomic number 13, and atomic weight 26.98. The isotope with mass number 27 is the only stable isotope. It is a soft, light, gray metal that resists corrosion when pure in spite of its chemical activity because of a thin surface layer of oxide. It is nonmagnetic and nonsparking. Its density is 2.6989 g/cm3, melting point 669.7°C and boiling point 1800°C. Its electrical resistivity is 2.824 μΩ-cm at 20°C, with temperature coefficient 0.0039°C-1, the same as copper's. Its thermal conductivity is 2.37 W/cm-K at 300K, and the linear coefficient of expansion is 23.86 x 10-6°C-1. The specific heat is 0.2259 cal/g-K, and the heat of fusion is 93 cal/g. The first ionization potential is 5.96V, second 18.74V and third 28.31V. Its electrode potential is 1.67V positive with respect to hydrogen. When near its melting point, it becomes "hot short" and crumbles easily. As a pure metal, it is quite soft, and must be strengthened by alloying with Cu, Mg, Si or Mn before it can be used structurally. Aluminium bronze is 90 Cu, 10 Al, a strong, golden-yellow alloy with excellent physical properties. The Young's modulus of pure aluminium is 10 x 106 psi, the shear modulus 3.8 x 106 psi, Poisson's ratio 0.33, and the ultimate tensile strength 10,000 psi, with 60% elongation. Pure aluminium is very ductile and malleable, and unsuitable as a structural material. Its hardness is 15 Brinell (500 kg, 10 mm). The useful wrought alloys contain 1-7% magnesium and 1% manganese. Its crystal form is face-centered cubic, with lattice constant a = 0.404 nm, and nearest-neighbor spacing of 0.286 nm.
The familiar strong aluminium alloy Duralumin should really be Düralumin, since it was originally the product of the Dürener Metallwerke in Germany. Düren is about halfway between Köln and Aachen in northwestern Germany. Dr Alfred Wilm was testing aluminium alloys there about 1909, and was surprised that an alloy 96 Al, 3.5 Cu, 0.5 Mg, which was not too impressive on a first test, strengthened greatly with a few day's rest after casting. What was happening was that all the Cu was in solution at 500°C, but at room temperature the solid could hold only 0.5% in solid solution. The hard intermetallic CuAl2 was forming slowly, making hard bits that would hinder the propagation of slip dislocations. This "age hardening" was the start of the general process of precipitation hardening that has been very important in strong alloys. The alloy was used in Zeppelins, and now is a major component of most aircraft, in the form of Alclad, in which the Duralumin is given a sheathing of pure aluminium to make it corrosion-resistant.
Liquid aluminium easily absorbs gases from the air, and these gases are expelled on solidification, causing flaws in castings. Casting alloys include silicon, and perhaps a little copper or nickel to help to avoid this. A eutectic is formed with aluminium at 11% silicon. Also, large crystals may be a problem if the aluminium is poured while too hot. The shrinkage on solidification is 0.2031 inches per foot for pure aluminium, 0.156 inches per foot for casting alloys.
The electron configuration of aluminium is 1s22s22p63s23p. The outer three electrons occupy three s2p hybrid orbitals that point in orthogonal directions. These electrons easily form covalent bonds, as in anhydrous AlCl3. This compound easily sublimates, showing that it is not ionic, and is partially hydrolyzed by water to release HCl gas. It cannot be formed by heating the hydrated form to drive off water. This only produces the oxide and HCl gas. It is now made commercially by heating aluminum oxide with carbon and chlorine. It is used in the refining of motor oil, and as a catalyst. Hydrated aluminium chloride is used as a personal deodorant. The acid environment it creates is unpleasant for microbes and mild enough to be non-irritating.
The spectroscopic ground state is 3s23p2 3P. The resonance line is at 396.15 nm, so the aluminium atom is not excited in the flame, and gives it no color. When the atom is excited, most of the lines are in the red or infrared. Aluminium is in column IIIA of the periodic table, which includes boron, aluminium, gallium, indium and thallium. Aluminium is the only common element in the group, and is considerably different from the others in physical and chemical properties. Boron is an acidic nonmetal, while gallium, indium and thallium are typical basic metals.
Aluminium should displace hydrogen from water because of its positive oxidation potential, but does not normally do so because of the protection by a surface layer of oxide. This oxide has the same density as the metallic aluminium, so it does not crack or wrinkle when it is formed, a lucky thing. A little HCl or NaOH that dissolves the oxide will permit the evolution of hydrogen. Aluminium pots and pans should not be used with acidic or strongly alkaline foods, that will dissolve the protective layer and allow the metal to be attacked. Aluminium is attacked by hydrochloric acid, but not by oxidizing acids like nitric. Aluminium tank cars are even used to transport nitric acid. If aluminium is amalgamated with mercury, the protective oxide layer is removed, the metal becomes very reactive, and will displace hydrogen from water. It can be used for handling cold nitric acid, but dissolves readily in alkalis to form aluminates, an example of the amphoteric behavior of aluminium. Formation of the oxide also prevents aluminium from being soldered with ordinary Sn-Pb solder, since the solder will not wet it. Sn-Zn solders can be used with aluminium. A suitable alloy is 85 Sn, 15 Zn, 5 Al. 60 Zn, 40 Cd has better corrosion resistance, but is harder to melt. The solder itself serves as a flux. With a suitable flux to protect the metal, aluminium can be welded.
Aluminium was first isolated by Hans Christian Oersted (1777-1851) in 1824 by reducing it from its oxide with potassium amalgam. The reaction is AlCl3 + 3K → 3KCl + Al. Two years later, Wöhler made the metal the same way. For some reason, Wöhler is usually considered the discoverer and got most of the fame. This Oersted is the same who discovered the magnetic effect of the electric current in 1820. In the Dictionary of Scientific Biography, the article on Oersted does not mention his work with aluminium. Refractory oxides such as alumina, silica or magnesia, were considered elements earlier, when all attempts to decompose them had failed. Only electrolysis made the decomposition possible, either directly or through the production of reactive metals like sodium and potassium.
The thermite reaction, discovered by Goldschmidt, is also a displacement reaction, but here aluminium reduces iron. The reaction is Fe2O3 + 2Al → 2Fe + Al2O3, which liberates a good deal of heat. The liquid metal produced is at about 2300°C, which is very hot. Powdered aluminium and rust in the approximate ratio of 1:3 are packed in a refractory crucible with a magnesium ribbon, or a powder of magnesium and barium peroxide, to ignite it. Either the red or black iron oxide can be used, giving "red Thermit" or "black Thermit." A trade name for the powder is Thermit. The vigorous reaction makes liquid iron or steel, which flows out of a hole in the bottom of the crucible into the mold and can be used for welding. The stock to be welded is usually preheated with a gas flame playing through the mold. The metal produced is about half the weight of the original mixture. This reaction is also called aluminothermic, and can be used for reduction of other metals, such as nickel, manganese or chromium.
Alumina, Al2O3, is the refractory oxide of aluminium, which does not melt below 2000°C. Corundum is the natural form of alumina, a very hard (9 on the Mohs scale, just below diamond) and heavy (4.02 g/cm3) that is a valuable gem when transparent. A chromium impurity makes ruby, while iron, cobalt or titanium makes sapphire. Fused alumina is called alundum, and is an artificial substitute for corundum, especially as an abrasive. Artificial rubies and sapphires of high quality are rather easily made. Artificial ruby is used in the ruby laser. The colors are not due to aluminium, but to the isolated impurity atoms. Emery is a natural corundum substance with magnetite or hematite as an impurity, which turns it black. Emery is an excellent abrasive, but should not be used where its electrical conductivity or magnetic impurities are deleterious. At 1700°C, alumina crystals become plastic and can be bent into any desired shape, for things like thread guides and phonograph needles.
Aluminium can be protected by a thick layer of oxide made by electrolysis, a process called "anodizing." The surface is first thoroughly cleaned and degreased in trichloroethylene, or by electrolysis, where the aluminium is the cathode and the grease driven off by the evolved hydrogen. Then the metal is made the anode in a bath of dilute sulphuric acid, and electrolyzed at 100 A/m2 and cell voltage 15V for about thirty minutes. The evolved oxygen combines with the aluminium to form a layer .007 to .015 mm thick (the normal oxide thickness formed in air is 13 nm). The layer is porous, and can be dyed by dipping in dye solutions. Finally, it is "sealed" by boiling in water, perhaps after dipping in a 1% nickel acetate solution or a surface application of linseed oil or similar. This renders the surface impermeable and resistant to staining.
Closely related to alumina is the hydroxide, Al(OH)3, usually formed as a gelatinous precipitate when aluminum compounds are hydrolyzed in water. If water is driven out of this precipitate by heating, a light, foamy solid results called activated alumina that will absorb moisture and other things, and can be reactivated by heating. This hydroxide reacts with both acids and bases according to the formula H+ + AlO2- + H2O = Al(OH)3 = Al+++ + 3OH-. Adding an acid removes OH-, driving the reaction to the right, while adding a base removes H+, driving the reaction to the left. Since it can go either way, aluminum hydroxide is called amphoteric, and is an excellent example of the type.
If we consider other hydroxides of elements in the same row of the periodic table, we see that NaOH is a base, and can react only with acids, while P(OH)5 = H+ + H2(PO)4- + H2O is an acid, and reacts only with bases. Only the first stage of ionization is shown. Aluminium is halfway between these two. This behavior depends on the ionic charge and the ionic radius, as is explained in any text on inorganic chemistry.
Aluminium sulphate, Al2(SO4)3·18H2O is a very useful aluminum compound, made from the oxide and sulphuric acid. When moistened, it becomes acidic because of hydrolysis, as just described. It is also called pickle alum from its use in giving sourness to pickles. The acidity can be used to produce carbon dioxide if combined with sodium bicarbonate, as in baking soda bread. Hydrolysis gives the gelatinous hydroxide, which is useful as a dye mordant, as a flocculent filter, and as size for paper. Dyes adhere very poorly to cotton, for example, so to dye cotton it is soaked in aluminium sulphate solution. The hydroxide impregnates the fibres and clings tightly to them. A dye then is adsorbed by the hydroxide, forming a lake. The word "lake" is from "lac," an insect resin originally used in lacquer which also had the property of absorbing dyes and becoming brightly colored.
Aluminium sulphate is found in nature as the rare mineral Kalinite, which is soluble in water, so it is found only near volcanoes and such places where it is make by the action of sulphuric acid on clays, such as on Lipari, near Vesuvius, and a few places in Germany and South America. Alum, KAl(SO4)2·12H2O (or twice this) is not as scarce, but is still rare, found in Europe and Utah. This mineral is partly hydrated, crystallizes in the hexagonal system, and is insoluble. It dissolves in H2SO4, however. Before 1600, when strong acids were not available, alum could be made by roasting this mineral and then extracting with water.
In general, an alum is of the form M+M'+++(SO4)2·12H2O. The monovalent cation can be K, Na, Cs, Rb, NH4 or Tl, while the trivalent cation can be Al, Fe or Cr. All form characteristic octahedral crystals (the symmetry may actually be monoclinic, as in KAl alum, but is not far from cubic). Alums are astringent, or styptic, and check bleeding, aiding those with exceptionally sharp or exceptionally dull razors. Aluminium acetate is a very soluble compound of aluminium, and behaves like the sulphate in the presence of water, precipiatating the hydroxide. Alums are much more soluble in hot water than in cold, so that crystals are easily grown from a solution that has become supersaturated by cooling. In cool water, KAl alum is less soluble than KFe, so it is very easy to eliminate small amounts of Fe ion by fractional crystallization. This ability to get very pure alum is very useful in dyeing, where a little Fe ion spoils the colour.
Alums are used with sodium bicarbonate in fire exinguishers as well as in baking powder, and for the same reason, the production of CO2. Cloth soaked in alum is fire-resistant, but I have not yet been able to discover why this should be so. Alum has also been mentioned as a fire retardant for wood, but again I do not know why (yet). These uses were probably unknown in ancient times, when alum was rare and expensive, used only as a styptic and in medicines, and perhaps rarely as a mordant. The famous mention of alum in Herodotus is probably not alum, as I maintain in another page on this site. Egypt had no alum deposits, and had no way to make the substance. Also, a shipload of this rather rare substance would be astonishing. The rebuilders of the Delphic temple would be at a loss over what to do with it.
Aluminium is not a very colorful element. It gives no coloration to the flame, and its compounds are relentlessly white. For a long time, the usual test for it consisted of forming Thénard's Blue, cobalt aluminate. Other ions interfere with this test (also giving blue) so the modern alternative of using Aluminon, a red dye, is now preferred. Aluminon is ammonium aurin tricarboxylate, whose structure is shown in the diagram. The chromophore is the region of alternating single and double bonds in the center. This dye will not adsorb on hydroxides of chromium, zinc, lead, tin or antimony to form a lake, having a decided preference for aluminium hydroxide. Aluminium dissolves slowly in dilute hydrochloric acid to make a clear solution. When ammonium hydroxide is added, a characteristic translucent gel precipitates, with a bluish tinge. Since aluminium is amphoteric, this gel will dissove in an excess of alkali. Other amphoteric cations can make a similar precipitate, so a test is necessary to confirm that aluminium is present. If a little aluminon is added, it makes a bright red lake with the gelatinous hydroxide, confirming the presence of aluminium.
Lapis lazuli, the mineral lazurite, is Na4-5Al3(SiO4)3S. This beautiful dark blue stone was greatly prized in antiquity, and is one aluminium compound that is not white. The best comes from northeastern Afghanistan. Finely ground, it made the pigment ultramarine. Ultramarine is now artificially made by fusing clay, carbon and sodium sulphate. Also not white is turquoise, Al2(OH)3PO4·H2O. These two stones are beautiful enough to compensate for all the white powders. Garnet and jade (jadeite, not nephrite) also contain some aluminium.
Aluminium salts are not poisonous, even when soluble.
Aluminum makes up 8.1% of the earth's crust by weight, the most common element after oxygen and silicon. It is probably restricted to the crust, where it is found in continental rocks like granite, whose principal minerals are silica, feldspar and mica, the last two containing important amounts of aluminium. Feldspar, KAlSi3O8, hydrates on weathering to form clay, such as kaolin, H2Al2(SiO4)2·H2O. Alunite, K2(SO)4·Al2(SO4)3 is a natural source of aluminium sulphate, resulting from the weathering of other igneous rocks.
From the time of its discovery until 1886, aluminium was produced by reduction with alkali metals, and was very expensive. The process was improved somewhat, but the price of aluminium in 1855 was $113 per pound, which precluded its use for anything but ornamental processes. The statue of Anteros (commonly supposed to be Eros) in Piccadilly Circus in London is an aluminium statue made with expensive aluminium. The Washington Monument in Washington, DC was capped with a 100-ounce aluminium pyramid costing $225. Napoleon III used aluminium tableware at his state dinners, and gave aluminium trinkets as gifts. No electrolytic process was available, because fusing aluminium oxide was much too difficult a task, even with electricity.
In 1886, two young men, C. M. Hall and Paul Héroult (both 1863-1914), working independently, came up with the solution. The melting point of alumina could be reduced to merely a red heat by adding it to fused cryolite, a rare mineral from Ivigtut in southwestern Greenland that dissolved alumina when fused. Fortunately, the cryolite was not consumed and could be re-used. The cryolite was held in a graphite-lined iron "pot" that was the cathode of an electrolytic cell. The anodes were carbon blocks hanging from the positive bus lowered into the molten cryolite from above. The passage of the current kept the cryolite hot and molten, at about 900°C, as well as reducing the aluminium to liquid metal, which collected on the bottom of the tank and could be drawn off periodically, as alumina was thrown in the top. The standard pig weighed 52 lb. The cathode reaction is Al+++ + 3e- → Al, and the anode reaction is C + 2O-- → CO2 + 4e-. The votage is about 5V. The alumina concentration began at about 5%, and when it decreased to 2% was replenished. Each pound of aluminium produced used 0.7 pound of carbon anode, and 0.1 pound of cryolite. Theoretically, a pound of aluminum should require about 5 kWH, but the actual amount is closer to 12.5-13 kWh per pound. About 100 pots may be connected in series, each pot requiring 6V at 8000-20,000A. The Hall-Héroult process uses huge amounts of electricity, so the cost of electricity is central to the economics of the process. Commercial production began in 1889 in Pittsburgh.
The Soderburg anode makes itself as it is consumed in the furnace. The carbon pitch-tar mixture is loaded at the top, and is baked by the furnace heat. This means that the process can continue indefinitely, without having to shut down to change anodes, an economic advantage. About half a ton of anode is required for every ton of aluminium produced, and 15 MWh of electricity. Another figure is 10 kWh per pound. If you do the arithmetic with what you pay for electricity at home, it turns out that aluminium would cost at least $0.50 per pound for electricity alone. In the U.S., electricity for aluminium reduction has been heavily subsidized to make the industry viable.
Since there was originally no convenient method for purifying the metal produced, it was necessary for the raw material, the alumina, to be purified instead. The Bayer process dissolved the impure alumina, usually in the form of bauxite, in molten sodium hydroxide. Since aluminium is amphoteric, as explained above, it dissolved but the iron and titanium hydroxides from the impurities did not. The purified alumina could then be recovered from the NaOH, and used in the Hall-Héroult process. Incidentally, this Hall is C. M. Hall, while the Hall of the Hall effect in electronics is E. H. Hall, who discovered his effect in 1879. Just to make things more confusing, there is also an H. E. Hall who worked much more recently in cryogenics. There is now the electrolytic Hoopes process for refining aluminium, so purity of the virgin aluminium is not as important as it once was. There are efforts to devise a process for the reduction of aluminium that is cheaper than the Hall-Héroult process and uses less electricity.
Although aluminium is very abundant in the earth's crust, compounds suitable for reduction to the metal are not. Silicates are not suitable, so attention must fall on the oxide, alumina. The principal source of alumina is the variable rock bauxite, named after the deposit at Les Baux-de-Provence, in southern France near Arles. It is a mixture of the mineral gibbsite, Al(OH)3, with lesser amounts of the denser and harder boemite and diaspore, AlO(OH). It is produced by lateric supergene weathering of clays in most cases, and is pisolitic in form. This is a gaudy way of saying that it is weathered by water from above, and consists of pea-like granules. A notable deposit is in Trinidad, but Australia, Brazil and Guinea have most of the resources. A famous deposit in the United States was in Arkansas, and other deposits were in Alabama and Georgia. Although the U.S. produces about 7,500,000 metric tons of aluminium yearly, half primary and half secondary (recovered from scrap), all of the bauxite is now imported. Alumina can also be made from alunite, mentioned above and found in the western U.S., but this source is not used.
Cryolite is sodium fluoaluminate, Na3AlF6. The density of cryolite is 2.90, and its melting point is 1000°C. It is a soft rock, with hardness only 2.5, but its curious distinction is an index of refraction of 1.364, very close to that of water. Because of this, particles of it in water seem to disappear. Its name, based on Greek, means "ice-rock," because of its appearance as a cloudy warm ice. Its very restricted occurrence was a great disadvantage to aluminium producers. A substitute made by fusing NaF, AlF3 and CaF2 now is used instead of the natural product. The molten bath in the pot is 59% AlF3, 21% NaF and 20% CaF2, with from 2% to 5% alumina. The fluorite was also used with natural cryolite.
The principal uses of aluminium are for aircraft, machinery, electrical conductors, and cooking utensils. Aluminium has largely created its own uses, and has not replaced steel in its traditional uses, as was once anticipated. The two principal structural metals are complementary more than competitive. Aluminium is very easy to work with in the shop, and has indeed replaced the more expensive brass in many applications. Aluminium, however, cannot be soldered, which is a severe disadvantage. It is very easy to make aluminium castings. At one time, electronic chassis were almost always aluminium, but printed-circuit boards have rendered metal chassis obsolescent.
Before electrolytic reduction, the expensive aluminium was only used for decorative items. We have mentioned the Eros statue in London, and the cap of the Washington Monument above, and some jewellery was also made. After electrolytic reduction appeared, the use of aluminium burgeoned. The first large application was to cooking utensils, which the ease of casting encouraged.
I have a cast-aluminium lever-type orange squeezer from the 1930's, marked "Wear-Ever," "Aluminum," "T.A.C.U.CO." and "Alcoa." It is still as good as new, works excellently, and is handsome.
Aluminium pots can be safely used with acid foods such as tomatoes, rhubarb, and fruits, since the acid environment strengthens the protective oxide layer. However, alkaline foods attack the aluminium, and should not be cooked in aluminium utensils. Hominy is probably an alkaline food, but there are not many. An old household hint is to boil vinegar in an aluminium pot to remove stains.
Aluminium is used for making small containers. Aluminium cans compete strongly with tinplate cans, especially for soft drinks. The aluminium can is made from two pieces only, the body deep-drawn and the top, of aluminium or steel, crimped on. They can be easily opened with pull tabs, that tear the metal along scored lines. All-aluminium cans are easily recycled. It is only necessary to remove the paint, and the high-quality scrap is ready for melting. The recycling of cans, tinplate as well as aluminium, is an easy and profitable process that should be greatly encouraged. Unlike much "recycling" it is actually an important saving.
The advantages of using aluminium as an electrical conductor were realized at an early date. This was especially true in high-voltage transmission, where the lightness of aluminum was a great convenience, and the larger diameter of the conductors was no drawback. Most high-voltage transmission lines are now aluminium, with steel cores to handle the tensile stress. Wiring codes allow copper, copper-clad aluminium and aluminium conductors. The NEC uses density 2.71 g/cm3, resistivity 2.790 μΩ-cm, tensile strength 19.0 kg/mm2 hard, 12.0 kg/mm2 annealed. Unless marked AL-CU, a receptacle may be used with copper or copper-clad conductors only. To get the same resistance per unit length, the AWG number of the aluminium conductor should be smaller than the AWG number of a copper conductor by 2. That is, #18 AWG copper is equivalent to #12 AWG aluminium.
Aluminium foil has replaced tinfoil for use in the kitchen and for packaging products. It is much cheaper, while resisting chemical attack just as well, and is equally safe around foods. Pure aluminium can be rolled into sheets as thin as .00127 mm.
Aluminium is used in making capacitors. Since it can be rolled into very thin foils, it was used in "paper" capacitors that were a roll of sheets of foil separated by thin paper dielectrics. These capacitors have now been superseded by plastic dielectrics and even thinner metal coatings. If some alkaline paste is sandwiched between two aluminium foils, and it is rolled up with an insulating sheet, a capacitor with a very thin dielectric of aluminium oxide is formed that is maintained by electrolysis. Some leakage is necessary to maintain the oxide, but it is very small, and the electrolytic capacitors that result are very economical of volume. Tantalum electrolytics are even more compact and have less leakage, but are more expensive.
Aluminium flakes can be used as a pigment in paint. This makes an excellent anti-corrosion paint for iron and other metals. It is also a good primer for wood. Its shiny, metallic appearance is also quite attractive. The flakes arrange themselves horizontally at the surface of the paint, separated by the vehicle, making a coherent, nonconducting layer that excludes ultraviolet light and moisture excellently. Little oil is needed to disperse the flakes, so the paint film is composed mainly of metal.
Aluminium surfaces make excellent mirrors, with a greater extent of spectral reflectivity than silver. Most mirrors are second-surface, where the metal is protected by the glass. The reflecting metal is most likely aluminium. Silver, and tin amalgam, were used earlier, but they are not as durable. Large reflecting telescopes have first-surface mirrors, and aluminium is invariably used these days. The mirror of the Hale 200-inch telescope is realuminized every five years.
All these uses contributed to a great increase in the use of aluminium even before the Second World War. Aluminium was the preferred structural material for aircraft after wood and fabric were superseded, and all sides in the war foresaw the need for increased aluminium capacity. The United States had a severe problem with bauxite supply that every means was taken to solve. New reduction plants were quickly built, mostly with government money, and were then leased to private companies for operation. New producers, such as Kaiser Aluminum and Reynolds Aluminum, joined the established Aluminum Corporation of America (Alcoa) to operate the government plants and profit by the war economy, which they did most handsomely. The aluminium demand was met, and in 1943 metal was even released for civilian uses, showing how successful the program had been.
W. N. Jones, Inorganic Chemistry (Philadelphia: Blakiston, 1949), Chapter 34.
J. L. Bray, Non-Ferrous Production Metallurgy (New York: John Wiley & Sons, 1947), Chapter 4.
R. A. Higgins, Engineering Metallurgy, 3rd ed. (London: The English Universities Press, 1971).
W. Alexander and A. Street, Metals in the Service of Man, 6th ed. (London: Pelican Books, 1976). Chapter 4.
Composed by J. B. Calvert
Created 7 November 2002
Last revised 9 January 2007