Iron

• 1. Introduction

  ›

  Iron sits in the periodic table as element 26, symbol Fe, a metal of the first transition series and group 8. By mass it is the most common element on the whole planet, forming much of Earth's outer and inner core. Yet on the surface it is rarely found in its pure metallic state, because it oxidizes so easily. In its native form it arrived mostly by meteorite, and the earliest iron tools were hammered from rocks that fell from the sky. To pull usable metal out of ordinary iron ores, humans had to build kilns and furnaces reaching 1500 degrees Celsius. That is about 500 degrees Celsius hotter than smelting copper demands. Why did a substance so abundant take so long to master? How did a metal that rusts into flakes become the backbone of modern industry and the iron at the center of every breath you take? The answers run from the cores of dying stars to the four grams of iron carried inside an adult human body.

• 2. Forged In Dying Stars

  ›

  56Fe is the most abundant iron isotope, and nuclear scientists prize it because it marks the most common endpoint of nucleosynthesis. Inside extremely massive stars, fusion chains run through the alpha process until they reach 56Ni, fourteen alpha particles, where conditions favor photodisintegration over building anything larger. That 56Ni has a half-life of about six days. Within the gas cloud of a supernova remnant, it decays by two successive positron emissions, first to radioactive 56Co and then to stable 56Fe. So iron becomes the most abundant element in the core of red giants and the most abundant metal in iron meteorites. Iron ranks as the sixth most abundant element in the universe and the most common refractory element. A tiny bit more energy could be wrung out by making 62Ni, which has a marginally higher binding energy, but stellar conditions do not allow it. Elements heavier than iron require a supernova, where rapid neutron capture starts from 56Fe nuclei. There is a stranger fate written into the far future. Assuming proton decay never happens, cold fusion by quantum tunnelling would slowly fuse light nuclei into 56Fe, while heavier nuclei would decay down toward it. Every star-sized object would eventually become a cold sphere of pure iron. Closer to home, the extinct radionuclide 60Fe, with a half-life of 2.6 million years, left its fingerprint in the meteorites Semarkona and Chervony Kut, evidence that it existed when the Solar System formed.

• 3. The Hidden Mantle

  ›

  Earth's inner and outer core together account for 35 percent of the entire planet's mass, and both are believed to be largely an iron alloy, possibly with nickel. Electric currents in the liquid outer core are thought to generate Earth's magnetic field. Mercury, Venus, Mars, and the Moon are all believed to carry metallic cores made mostly of iron, and the M-type asteroids may be partly or mostly metallic iron alloy. Ferropericlase, a solid solution of periclase and wuestite, makes up about 20 percent of the volume of Earth's lower mantle. That makes it the second most abundant mineral phase there, after silicate perovskite, and the major host for iron in the lower mantle. Silicate perovskite in its magnesium iron form may account for up to 93 percent of the lower mantle. It is considered the most abundant mineral in the Earth, making up 38 percent of the planet's volume. Despite all this, the iron in Earth's crust amounts to only about 5 percent of the crust's mass, making it merely the fourth most abundant element there, after oxygen, silicon, and aluminium. The rare iron meteorites remain the main form of natural metallic iron on the surface. The Inuit in Greenland have been reported to take iron from the Cape York meteorite for their tools and hunting weapons. About 1 in 20 meteorites consist of the iron-nickel minerals taenite and kamacite, the latter running as high as 90 to 95 percent iron.

• 4. Color, Rust, And Ruin

  ›

  Pristine pure iron surfaces shine a mirror-like silvery-gray, yet that surface rarely survives. Iron reacts readily with oxygen and water to make brown-to-black hydrated iron oxides, the substance everyone knows as rust. The cruelty of rust is geometric. Unlike oxides that form a protective passivating layer, rust takes up more volume than the metal beneath it, so it flakes away and exposes fresh surface to keep corroding. Materials full of finely ground iron oxides have served as yellow, red, and brown pigments since pre-historical times. They tint the Painted Hills in Oregon, the Buntsandstein, and through Bath stone they give many historical buildings their yellowish cast. The proverbial red of the Martian surface comes from an iron oxide-rich regolith. This appetite for corrosion carries a price measured against the whole world. Protecting iron from rust costs over 1 percent of the world's economy. People fight back with painting, galvanization, passivation, plastic coating, bluing, and cathodic protection, all aimed at shutting out water and oxygen. In cities the electrolyte driving the reaction is usually iron(II) sulfate, formed when atmospheric sulfur dioxide attacks the metal. By the sea, it is salt particles carried on the air.

• 5. From Sky-Metal To Steel

  ›

  Beads made from meteoric iron, dated to 3500 BC or earlier, were found at Gerzeh in Egypt by G. A. Wainwright. Their 7.5 percent nickel content is a signature of meteoric origin, since crustal iron carries only minuscule nickel impurities. A dagger of meteoric iron lay in the tomb of Tutankhamun, its iron, cobalt, and nickel proportions matching a meteorite found in the same area. Smelting changed everything, but slowly. Samples of smelted iron from Asmar in Mesopotamia and Tall Chagar Bazaar in northern Syria date between 3000 and 2700 BC. The Hittites, who built an empire in north-central Anatolia around 1600 BC, appear to be the first to understand producing iron from its ores. They began smelting iron between 1500 and 1200 BC, and after their empire fell in 1180 BC the practice spread across the Near East, opening the Iron Age. Cast iron was first produced in China during the 5th century BC, with the earliest artifacts found in what is now Luhe County, Jiangsu. The technology kept advancing in Europe. In 1709 Abraham Darby I established a coke-fired blast furnace for cast iron, replacing charcoal. Cheap iron helped drive the Industrial Revolution and made structures like the first iron bridge of 1778 possible, a bridge that still stands. In 1783 Henry Cort patented the puddling process for refining iron. Then, in the late 1850s, Henry Bessemer invented a process that blew air through molten pig iron to make mild steel cheaply, ending large-scale wrought iron production. So central did the metal become that many languages call the railway iron road, from the French chemin de fer to the German Eisenbahn.

• 6. A Metal Of Many Faces

  ›

  Three arrangements of iron atoms appear in solid iron at ordinary pressures, denoted alpha, gamma, and delta. As molten iron cools past its freezing point of 1538 degrees Celsius, it forms the delta allotrope with a body-centered cubic structure. Cooling to 1394 degrees Celsius turns it into face-centered cubic gamma-iron, known as austenite. At 912 degrees Celsius and below it returns to body-centered cubic alpha-iron. Magnetism is the metal's most famous trick. Below its Curie point of 770 degrees Celsius, alpha-iron switches from paramagnetic to ferromagnetic, as the spins of two unpaired electrons in each atom align with their neighbors. Without an external field, the atoms partition into magnetic domains about 10 micrometers across, often pointing different directions, so an ordinary piece of iron carries almost no net field. Apply a field and the aligned domains grow at the expense of others. Pin them with impurities or grain boundaries and the effect survives, turning the object into a permanent magnet. Naturally magnetized pieces of magnetite, called lodestones, gave navigation its earliest compasses. Iron's variety extends to chemistry, where it forms compounds across oxidation states from minus 2 to plus 7. It cannot reach the group oxidation state of plus 8 that its heavier relatives ruthenium and osmium achieve. Iron has four stable isotopes, with 56Fe making up 91.754 percent of natural iron. One quirk of behavior keeps mercury in iron flasks: iron does not form amalgams with mercury, so the liquid metal is traded in standardized 76 pound flasks made of iron.

• 7. Ferrocene And The Chemists

  ›

  1951 brought a landmark to organometallic chemistry, the discovery of the remarkably stable sandwich compound ferrocene, by Pauson and Kealy and independently by Miller and colleagues. Its surprising molecular structure was worked out only a year later by Woodward and Wilkinson and by Fischer. That single iron compound transformed the field in the 1950s, and ferrocene remains one of its most important tools and models. Prussian blue, an old iron-cyanide complex, has long served as a pigment and provides the traditional blue in blueprints. Its formation doubles as a simple wet chemistry test that distinguishes Fe2+ from Fe3+ solutions. Iron pentacarbonyl, in which a neutral iron atom binds five carbon monoxide molecules, can be made into carbonyl iron powder, a highly reactive form of metallic iron. Iron earns its place as a workhorse partly because it is inexpensive and nontoxic, which has driven much effort into iron-based catalysts and reagents. Iron catalysts traditionally run the Haber-Bosch process for ammonia and the Fischer-Tropsch process that turns carbon monoxide into hydrocarbons for fuels and lubricants. Even older chemistry leaned on it. In 1774 Antoine Lavoisier reacted water steam with metallic iron inside an incandescent iron tube to produce hydrogen, work that helped demonstrate the conservation of mass and pushed chemistry from a qualitative science toward a quantitative one.

• 8. The Iron Inside Us

  ›

  About four grams of iron, roughly 0.005 percent of body weight, sit inside an adult human, three quarters of it in hemoglobin. Only about one milligram is absorbed each day, because the body recycles its hemoglobin for the iron it already holds. The body has no regulated way to excrete iron, so it controls levels almost entirely by regulating uptake. In hemoglobin the iron occupies one of four heme groups with six coordination sites. Four are held by nitrogen atoms in a porphyrin ring, the fifth by a histidine residue, and the sixth reserved for oxygen. When deoxyhemoglobin has no oxygen, the high-spin Fe2+ ion sits about 55 picometers above the ring, too large to fit inside it. Picking up oxygen flips the ion to low-spin, shrinks its radius about 20 percent, and lets it drop into the now-planar ring. That shift ripples through the protein, raising every subunit's affinity for more oxygen. Carbon monoxide is poisonous because it binds hemoglobin like oxygen but far more strongly, blocking transport; the bound form is called carboxyhemoglobin. Iron is the most common nutritional deficiency in the world, and untreated it leads to iron-deficiency anemia, marked by too few red blood cells and too little hemoglobin. Excess is dangerous too. A genetic defect mapping to the HLA-H gene region on chromosome 6 lowers hepcidin and can cause iron overload, or hemochromatosis, estimated to cause 0.3 to 0.8 percent of all metabolic diseases of Caucasians. For that reason people should not take iron supplements unless they are deficient and have consulted a doctor.