The allure of aluminium | Nature Chemistry

Daniel Rabinovich outlines the history, properties and uses of aluminium — one of the most versatile, pervasive and inexpensive metals today, yet it was considered a rare and costly element only 150 years ago.

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It is hard to believe that aluminium was once more expensive than gold and that, in the mid-nineteenth century, Napoleon III used silverware made of the light metal when he really wanted to impress his guests at stately dinners. Even though element 13 is the most abundant metal in the Earth's crust (∼8%) and is present in more than 270 different minerals, its high affinity for oxygen and the chemical stability of its oxides and silicates precluded its isolation in pure form for a long time. The first pure sample of aluminium was obtained in 1827 by the German chemist Friedrich Wöhler, who also began studying its fascinating physical and chemical attributes.

Credit: QUILLIVIC © LA POSTE, 1986

The French chemist Henri Sainte-Claire Deville (1818–1881) developed a method of preparing larger quantities of aluminium in 1854, and soon published the first comprehensive book describing its manufacture, properties and emerging applications1.

The attractive properties of the newfangled metal quickly became clear, including low density, high tensile strength and malleability, good thermal and electrical conductivity, and a remarkable resistance to corrosion. Jules Verne eloquently wrote in From the Earth to the Moon (1865) that “This valuable metal possesses the whiteness of silver, the indestructibility of gold, the tenacity of iron, the fusibility of copper, the lightness of glass. It is easily wrought, it is very widely distributed, forming the basis of most of the rocks, is three times lighter than iron, and seems to have been created with the express purpose of furnishing us with the material for our projectile.” The price of aluminium, however, was still comparable to that of silver, which hampered the development of large-scale applications and motivated the search for an alternative and more economical preparation process.

It was only in 1886 that Charles M. Hall in the US and Paul L. T. Héroult in France, almost simultaneously and completely independently, devised aluminium production processes that relied on the electrolysis of alumina (Al2O3) dissolved in molten cryolite (Na3AlF6). An efficient process for the extraction and purification of alumina from bauxite, the most important aluminium ore, was developed within a couple of years by the Austrian chemist Karl Josef Bayer, son of the founder of the famous German chemical and pharmaceutical company, and the 'Hall–Héroult' process became economically viable. By the early 1960s element 13 became the most widely used non-ferrous metal in the world, even more so than copper.

Applications of aluminium and its alloys range from construction and the transportation industry to the manufacture of electric power lines, packaging materials, cooking utensils and a myriad of other household goods. Another important feature of this ubiquitous metal, one that has significant economic and environmental consequences, is the ease with which it can be recycled. The recovery of secondary aluminium requires only about 5% of the energy necessary to produce new metal from bauxite, while also leading to a decreased use of landfill space and a reduced emission of greenhouse gases.

In contrast to the relatively short history of the pure metal, compounds of aluminium have long been known: alum, a hydrated sulfate of potassium and aluminium, KAl(SO4)2·12H2O, was used as an astringent and a dyeing mordant in ancient Greece and Rome. Aluminium chloride, AlCl3, a common Lewis acid, is extensively applied in Friedel–Crafts acylation and alkylation reactions, and aluminium chlorohydrate, Al2Cl(OH)5, is the active ingredient in many antiperspirants. Large quantities of methylaluminoxane, a generic name used to describe the ill-defined mixture of species obtained by partial hydrolysis of trimethylaluminium, are employed in the Ziegler–Natta polymerization of olefins.

The availability of an ever-increasing variety of aluminium coordination complexes has also prompted many recent developments in the chemistry of this metal, often with potential applications to catalysis and organic synthesis2. Other active areas of research range from the preparation of unusual aluminium(I) compounds, including organometallic species3 and metalloid clusters4, to the synthesis of Schiff base derivatives that effectively break down organophosphate nerve agents and pesticides5.

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The element once dubbed the magic metal by National Geographic continues to be a source of inspiration for scientists, engineers and even artists and designers6. Let us remember its rich chemistry, fascinating history and multifarious applications the next time we wrap a sandwich in aluminium foil or drink a carbonated beverage from a can!

Hall’s Early Experiments with Aluminum

Hall took his first formal course in chemistry as a junior in college. Earlier, with Jewett's guidance and encouragement, he had worked on aluminum chemistry and other projects in Jewett's laboratory and in his own laboratory at home. He heard Jewett lecture on the chemistry of aluminum, display his sample of the metal, and predict the fortune that awaited the person who devised an economical method for winning aluminum from its oxide ore. To a fellow student, Hall declared that he intended to be that person.

After many unsuccessful experiments with chemical methods of reducing aluminum ores to the metal, Jewett and Hall turned to electric current to provide the powerful reducing conditions that were needed. To obtain electricity in a college town in the 1880s, one had to assemble batteries. Hall and Jewett used Bunsen Grove cells, which consist of a large zinc metal electrode in a sulfuric acid solution that surrounds a porous ceramic cup containing a carbon rod immersed in concentrated nitric acid. Assembling enough of these cells to provide sufficient electrical energy for aluminum production was a large undertaking. The eventual laboratory process used about one pound of zinc electrodes, hand cast by Hall, to obtain one ounce of aluminum.

Hall did the first experiments with electricity in Jewett's laboratory during his senior year of 1884/85. He prepared aluminum fluoride from hazardous hydrofluoric acid in special lead vessels, and he passed a current through aluminum fluoride dissolved in water. Unfortunately, this system produced only unwanted hydrogen gas and aluminum hydroxide at the negative electrode.

After graduation, Hall continued the work in the woodshed behind his family's house. He experimented with molten fluoride salts as water-free solvents. He knew that the fluoride salts had the advantage over previously studied chloride salts of not absorbing water from the air. Hall was aware of Richard Grïtzel's success in obtaining magnesium metal by using an electric current in a magnesium chloride melt as reported in the Scientific American in 1885.

To work with molten fluoride salts, he needed a furnace capable of producing and sustaining higher temperatures than the coal-fired furnace of his earlier experiments. For this purpose, Hall adapted a second-hand, gasoline-fired stove to heat the interior of a clay-lined iron tube. Despite the higher temperature of this furnace, he was unable to melt calcium, aluminum or magnesium fluorides. Potassium and sodium fluorides melted but did not dissolve useful amounts of aluminum oxide.

Hall moved on to experiment with cryolite (sodium aluminum fluoride) as a solvent. He made cryolite, found that it would melt in his furnace, and showed that it would dissolve more than 25% by weight of aluminum oxide. The melting point of cryolite is 1000° C, an exceptionally high temperature for electrochemistry. He did this crucial experiment early in February 1886 and repeated it the next day for his sister Julia to witness.

Six days later, Hall first attempted to prepare aluminum metal by passing an electric current through a solution of aluminum oxide in molten cryolite. He immersed graphite rod electrodes into the fiery solution in a clay crucible and let the current run for a while. In Julia's presence, he poured the melt into a frying pan and broke apart the cooled mass but found no aluminum. There was only a grayish deposit on the negative electrode, a deposit that did not have the shiny metallic appearance of aluminum. After repeating this process several times, Hall realized that this deposit was probably silicon from silicates dissolved out of the clay crucible. If he had not been acquainted with the appearance of metallic aluminum from seeing Jewett's sample, Hall might have been slower to interpret this false result.

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