Iron - chemistry.

Iron - chemistry. I INTRODUCTION Iron, symbol Fe (Latin ferrum, "iron"), magnetic, malleable, silvery white metallic element. The atomic number of iron is 26; iron is one of the transition elements of the periodic table (see Periodic Law). Metallic iron was known and used for ornamental purposes and weapons in prehistoric ages; the earliest specimen still extant, a group of oxidized iron beads found in Egypt, dates from about 4000 BC. The archaeological term Iron Age properly applies only to the period when iron was used extensively for utilitarian purposes, as in tools, as well as for ornamentation (see Metalwork). The beginnings of modern processing of iron can be traced back to central Europe in the mid-14th century II BC. PROPERTIES Pure iron has a hardness that ranges from 4 to 5. It is soft, malleable, and ductile. Iron is easily magnetized at ordinary temperatures; it is difficult to magnetize when heated, and at about 790°C (about 1450°F) the magnetic property disappears. Pure iron melts at about 1538°C (about 2800°F), boils at 2861°C (5182°F), and has a specific gravity of 7.87. The atomic weight of iron is 55.845. The metal exists in three different forms: ordinary, or ?-iron (alpha-iron); g-iron (gamma-iron); and ?-iron (delta-iron). The internal arrangement of the atoms in the crystal lattice changes in the transition from one form to another. The transition from ?-iron to g-iron occurs at about 910°C (about 1700°F), and the transition from giron to ?-iron occurs at about 1400°C (about 2600°F). The different physical properties of all allotropic forms and the difference in the amount of carbon taken up by each of the forms play an important part in the formation, hardening, and tempering of steel. Chemically, iron is an active metal. It combines with the halogens (fluorine, chlorine, bromine, iodine, and astatine), sulfur, phosphorus, carbon, and silicon. It displaces hydrogen from most dilute acids. It burns in oxygen to form ferrosoferric oxide, Fe3O 4. When exposed to moist air, iron becomes corroded, forming a reddish-brown, flaky, hydrated ferric oxide commonly known as rust. The formation of rust is an electrochemical phenomenon in which the impurities present in iron form an electrical "couple" with the iron metal. A small current is set up, water from the atmosphere providing an electrolytic solution. Water and soluble electrolytes such as salt accelerate the reaction. In this process the iron metal is decomposed and reacts with oxygen in the air to form rust. The reaction proceeds faster in those places where rust accumulates, and the surface of the metal becomes pitted. See Corrosion. When iron is dipped into concentrated nitric acid, it forms a layer of oxide that renders it passive--that is, it does not react chemically with acids or other substances. The protective oxide layer is easily broken through by striking or jarring the metal, which then becomes active again. III OCCURRENCE Metallic iron occurs in the free state in only a few localities, notably western Greenland. It is found in meteorites, usually alloyed with nickel. In chemical compounds the metal is widely distributed and ranks fourth in abundance among all the elements in Earth's crust; next to aluminum it is the most abundant of all metals. The principal ore of iron is hematite, which is mined in the United States in Minnesota, Michigan, and Wisconsin. Other important ores are goethite, magnetite, siderite, and bog iron (limonite). Pyrite, the sulfide ore of iron, is not processed as an iron ore because it is too difficult to remove the sulfur. For details of the processing of iron ore, see Iron and Steel Manufacture. Small amounts of iron occur in combination in natural waters, in plants, and as a constituent of blood. IV USES Pure iron, prepared by the electrolysis of ferrous sulfate solution, has limited use. Commercial iron invariably contains small amounts of carbon and other impurities that alter its physical properties, which are considerably improved by the further addition of carbon and other alloying elements. By far the greatest amount of iron is used in processed forms, such as wrought iron, cast iron, and steel. Commercially pure iron is used for the production of galvanized sheet metal and of electromagnets. Iron compounds are employed for medicinal purposes in the treatment of anemia, when the amount of hemoglobin or the number of red blood corpuscles in the blood is lowered. Iron is also used in tonics. V PRODUCTION In the early 1990s, annual United States production of iron ore exceeded 56 million metric tons. In the same period, world production was nearly 920 million metric tons. The estimated worth of usable ore produced in 1990 in the United States was more than $1.7 billion. VI COMPOUNDS Iron forms ferrous compounds in which it has a valence of +2 and ferric compounds in which it has a valence of +3. Ferrous compounds are easily oxidized to ferric compounds. The most important ferrous compound is ferrous sulfate (FeSO4), called green vitriol or copperas; it usually occurs as pale-green crystals containing seven molecules of water of hydration. It is obtained in large quantities as a by-product in pickling iron and is used as a mordant in dyeing, as a tonic medicine, and in the manufacture of ink and pigments. Ferric oxide, an amorphous red powder, is obtained by treating ferric salts with a base or by oxidizing pyrite. It is used both as a pigment, known as either iron red or Venetian red; as a polishing abrasive, known as rouge; and as the magnetizable medium on magnetic tapes and disks. Ferric chloride, obtained as dark-green, lustrous crystals by heating iron in chlorine, is used in medicine as an alcoholic solution called tincture of iron. The ferrous and ferric ions combine with cyanides to form complex cyanide compounds. Ferric ferrocyanide (Fe4[Fe(CN)6]3), a dark-blue, amorphous solid formed by the reaction of potassium ferrocyanide with a ferric salt, is called Prussian blue. It is used as a pigment in paint and in laundry bluing to correct the yellowish tint left by the ferrous salts in water. Potassium ferricyanide (K3Fe(CN)6), called red prussiate of potash, is obtained from ferrous ferricyanide (Fe3[Fe(CN)6] 2; also called Turnbull's blue), and is used in processing blueprint paper. Iron also undergoes physiochemical reactions with carbon that are essential to the formation of steel. Microsoft ® Encarta ® 2009. © 1993-2008 Microsoft Corporation. All rights reserved.