Iron
Iron, (Latin ferrum,”iron”) symbol Fe, 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. The beginnings of modern processing of iron can be traced back to central Europe in the mid-14th century 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 1535° C (about 2795° F), boils at 2750° C (4982° F), and has a specific gravity of 7.86. The atomic weight of iron is 55.847.
The metal exists in three different forms: ordinary, or a-iron (alpha-iron); g-iron (gamma-iron); and δ-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 a-iron to g-iron occurs at about 910° C (about 1700° F), and the transition from g-iron to δ-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. 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.
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.
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. The beginnings of modern processing of iron can be traced back to central Europe in the mid-14th century 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 1535° C (about 2795° F), boils at 2750° C (4982° F), and has a specific gravity of 7.86. The atomic weight of iron is 55.847.
The metal exists in three different forms: ordinary, or a-iron (alpha-iron); g-iron (gamma-iron); and δ-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 a-iron to g-iron occurs at about 910° C (about 1700° F), and the transition from g-iron to δ-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. 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.
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.
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