IRON (AS. īren, īsern, Goth. eisarn, OHG. īsarn, īsan, Ger. Eisen, from OIr. iarn, Welsh haiarn, Corn. hoern, Bret. hoiarn, iron; possibly connected ultimately with Lat. æs, Goth. aiz, OHG. ēr, Ger. Erz, bronze, AS. ār, Eng. ore, Skt. ayas, metal). The most abundant and useful of the metals. Unlike many of the common metals, iron is found only rarely in the native state, its occurrence being limited to meteorites which fall upon the earth from extra-terrestrial regions. Combined with oxygen and other elements, however, it is widely distributed as a constituent of rocks, and forms extensive deposits in many parts of the world.
Early History. Articles of stone, bronze, and iron have been found together on the site of the Swiss lake dwellings, but the most ancient specimens of iron at present known come from Egyptian and Assyrian ruins. There is in the British Museum a piece of iron taken from the Pyramid of Gizeh, which is believed to date from about B.C. 4000, and also an axe-head of Egyptian manufacture dating from B.C. 1370. Though the earliest pieces of iron now known came from Egypt, it is thought that probably the Assyrians were the first to use the metal freely in the manufacture of tools, weapons, and ornaments. Numerous samples of iron implements dating from B.C. 800, and including the fragment of a saw 44 inches long, were obtained from the ruins of Nimrud. In India, the famous iron pillar at Kutub, near Delhi, standing 22 feet above ground and weighing about 6 tons, dates from B.C. 400. It was made by welding disks of metal together, indicating a remarkable degree of skill on the part of those early iron-workers.
The Greeks were familiar with the uses of iron at least six hundred years before the Christian Era, although the metal was very scarce. It was not until the Roman Empire was fairly established that the use of iron became general over civilized Europe. The Romans made both wrought iron and steel, using the latter metal for swords and other edged implements. The ancient Britons, at the time of the invasion of Britain by Julius Cæsar, B.C. 55, were found to use swords, spears, hooks, and scythes of iron, indicating a familiarity with the metal for a considerable period previous. During the Roman occupation iron manufacture was vigorously developed, and in the succeeding Saxon age it seems also to have flourished. After the Norman Conquest little attention was paid to iron manufacture in England. In Germany, however, the art flourished vigorously, iron implements being exported to other countries in considerable quantities. Until about the middle of the fourteenth century all iron was produced by the direct process, the product being wrought iron or steel, as the case might be, according to the details of the process. About 1350, however, cast iron began to be made in Germany, and the beginning of the modern process of iron and steel manufacture was inaugurated.
The succeeding history of iron and steel manufacture will be continued in the article following under Cast Iron, Wrought Iron, and Steel.
Iron (symbol, Fe; atomic weight, 56), in the chemically pure state, is a silver-white metal that crystallizes in the isometric system. Its specific gravity is 7.84, and its melting-point between 1400° and 1500° C. (2500° and 2700° F.). It is the most tenacious of all the ductile metals, and it may be rolled into sheets so thin that the weight of a sheet of given size will be less than the weight of a sheet of paper of the same size. The magnetic properties are well known. Pure iron may be prepared by the prolonged action of a weak current of electricity on a solution containing pure ferrous sulphate and sal ammoniac or sulphate of magnesium, and heating the precipitated metal with a view of freeing it from ‘occluded’ hydrogen and diminishing its brittleness. Another method of obtaining chemically pure iron consists in preparing pure ferric hydroxide by adding ammonia to the solution of some pure iron salt, and then heating the hydroxide in a stream of hydrogen gas. Iron and platinum are the only metals that may be soldered together immediately, i.e. without the use of any soldering material. With mercury iron refuses to combine directly, the iron amalgam that has been used for electrical machines being made by a somewhat complicated process with the aid of the amalgam of sodium. When immersed in fuming nitric acid, iron becomes coated with insoluble oxide, and as it then refuses to dissolve in acids (unless the coating be removed), it is said to be in a passive state. Dry air or oxygen gas has no effect on iron. The ‘rusting’ of iron in ordinary atmospheric air is due to the presence of moisture; the fact that the rust thus formed is invariably found to contain ammonia would seem to indicate that iron reacts chemically with the moisture of the air, combining with its oxygen and setting free its hydrogen, which, in the nascent state, forms ammonia with the nitrogen of the air. Rusting may accordingly be prevented by covering iron with a waterproof coating or some paint or varnish, or with a coating of some metal like lead, tin, copper, nickel, or preferably zinc.
The Oxides of Iron. With oxygen, iron forms three distinct compounds: Ferrous oxide, FeO; ferric oxide, Fe2O3; and ferric anhydride, FeO3, or Fe2O6. When oxidized at high temperatures, iron yields the substance Fe3O4; but this is really a combination of ferrous and ferric oxides. Ferrous oxide, or rather ferrous hydroxide, Fe(OH)2, may be obtained by adding caustic soda to a solution of green vitriol (ferrous sulphate). It is a white compound readily absorbing oxygen from the air, even if kept under water, and thus changing into ferric oxide; the oxidation causes its color to change gradually from white to green, gray, and brown. It is but sparingly soluble in water, the solution having an alkaline reaction. If boiled with a solution of caustic potash, ferrous hydroxide attacks the water of the solution, setting free its hydrogen and combining with its oxygen to form ferric hydroxide. Ferric oxide, in its anhydrous form, is found extensively as hematite. It is prepared artificially, by heating green vitriol, for use as an oil paint for wood, being known as colcothar. Ordinary rouge, which is used for polishing glass and metals, is artificial ferric oxide reduced to a fine powder. Ferric oxide may be obtained in the form of crystals having a dark-violet color, by heating green vitriol with common salt. If strongly heated, ferric oxide loses its property of readily dissolving in acids, and can then be dissolved only in strong acids and only at a high temperature. If heated to a white heat, it loses part of its oxygen and becomes converted into Fe3O4, which exhibits marked magnetic properties. The hydrate of ferric oxide, as ordinarily obtained by adding alkalies to solutions of ferric salts, has the composition 2Fe2O3.3H2O, and is readily soluable in acids. Another hydrate of ferric oxide has the composition Fe2O3.H2O, and, like the anhydrous oxide, does not, if heated, readily dissolve in acids. By exactly neutralizing a solution of ferric chloride with alkali, and subjecting the resulting liquid to a process of hydrolysis, pure ferric oxide may be obtained in aqueous solutions. But, like other colloidal solutions (see Colloids), that of ferric oxide is unstable, and the oxide readily passes from the soluble to the insoluble form. The hydrated forms of the oxide may lie converted into the ordinary anhydrous form by the application of a moderate heat; at a certain point of this process, the substance suddenly becomes incandescent—showing that a peculiar molecular change is taking place in it—and after that it is found to have lost its property of readily dissolving in acids. The exact nature of the change is unknown. Ferric anhydride is unknown in the free state. It is the hypothetical anhydride of ferric acid. H2FeO4, which is likewise unknown in the free state, but certain of whose salts may be readily prepared. Thus, potassium ferrate, K2FeO4, may be obtained by heating small pieces of iron with the chlorate of potassium. From the existence of such salts it is evident that, much unlike ferrous oxide, which is distinctly alkaline in reaction, the peroxide of iron acts as a feeble acid. The tendency of peroxidized iron to pass into the stable ferric state is even greater than the tendency of ferrous iron to pass into that state, and hence the ferrates act as powerful oxidizers, readily burning such substances as oxalic acid, readily changing manganous oxide into manganese dioxide, etc.
The Salts of Iron. Corresponding to ferrous oxide and ferric oxide, respectively, are two series of iron salts, ferrous salts and ferric salts.
The action of acids on metallic iron, in the absence of oxidizing agents, causes the formation of ferrous salts, among which deserve mention the sulphate, the sulphide, the chloride, and the oxalate. Ferrous sulphate, known as green vitriol, or iron vitriol, is the substance from which the compounds of iron are generally prepared. The sulphate itself is obtained as a by-product in certain industrial processes, and may be prepared by the action of sulphuric acid on metallic iron. In its ordinary, hydrated form its composition corresponds to the formula FeSO47H2O. It has a greenish color that is scarcely perceptible when the salt is dissolved in water; the solution readily takes up oxygen, which causes the formation of ferric sulphate, and hence must be kept in sealed vessels out of contact with air, if it is to be preserved unchanged. Ferrous sulphate is used for a variety of purposes in the arts; it is employed in making fuming sulphuric acid, in dyeing, as a disinfectant, in making colcothar and rouge, etc. Ferrous sulphide, FeS, which has been found in many meteoric stones, may be made by heating iron filings with flowers of sulphur. It is largely used in chemical laboratories for the preparation of sulphureted hydrogen, which it yields on coming into contact with dilute sulphuric or hydrochloric acid. Ferrous chloride, or rather its hydrated form, FeCL2.4H2O, may be prepared by the action of hydrochloric acid on metallic iron. The crystalline anhydrous chloride, FeCl2, may be prepared by the action of gaseous hydrochloric acid on red-hot iron. Ferrous oxalate, which is a powerful reducing agent, is used as a developer in photography, potassium-ferrous oxalate being used for the same purpose. Another important compound containing iron in the ferrous state is the well-known potassium ferrocyanide, or yellow prussiate of potash, which may be found described under Hydroferrocyanic Acid.
Among the ferric salts deserve mention the chloride, the sulphide, the nitrate, and the phosphate. Ferric chloride, Fe2Cl6, is a volatile and extremely hygroscopic salt prepared by the action of chlorine upon red-hot metallic iron. Its solutions in water have a brown color which may possibly be due not to the ferric chloride itself, but to the formation of basic chlorides of iron, which are eventually precipitated out of solutions of ferric salts. Commercial ferric chloride contains a considerable percentage of water, and hence contains basic chlorides, probably some free ferric hydroxide, etc. It is prepared by dissolving ordinary ferric hydroxide in hydrochloric acid. Ferric sulphide, FeS2, occurs in nature abundantly as iron pyrite; it is used for the preparation of sulphurous anhydride in manufacturing sulphuric acid and in bleaching. Ferric nitrate, Fe2(NO3)6, is obtained by dissolving metallic iron in an excess of cold nitric acid and allowing the solution to evaporate in a vacuum; the crystals thus obtained correspond to the formula Fe2(NO3)6.9H2O, and melt at 35° C. In aqueous solution, the nitrate gradually decomposes unless an excess of free nitric acid is present. Ferric phosphate, FePO4, is an insoluble white substance formed when acid sodium phosphate is added to solutions of ferric acetate. Another compound containing iron in the ferric state, viz. potassium ferricyanide, may be found described under Hydroferricyanic Acid.
Medicinal Uses of Iron Compounds. Iron itself and a number of its compounds are used in medicine in the form of various preparations; in the stomach all such compounds are converted into ferric chloride, and to a small extent into ferrous chloride. One of the best medicinal compounds of iron is ferric chloride, the evil effects of whose strongly acid properties may be avoided by the addition of bicarbonate of sodium. Another way to avoid the undesirable effects of acid compounds of iron is to administer them in the form of coated pills which may pass through the stomach unchanged, the acidity being then neutralized in the alkaline juices of the intestine. The constipating effect of iron compounds is well known, but is generally somewhat exaggerated; this effect may be readily avoided by the use of suitable purgatives. To avoid indigestion, iron compounds should not be taken shortly before or after meals. In the mouth, iron salts may (if acid) attack the enamel of the teeth, and by combining with sulphur (from food or the tartar of the teeth) form a black deposit of ferrous sulphide on the teeth and the tongue. For these reasons iron preparations are usually administered through a glass tube, and the mouth is to be carefully rinsed immediately after taking the dose.
Besides constituting the best-known local astringents for external application, iron salts are extensively used as a remedy for many forms of anæmia and the conditions caused by them, the best results being obtained by the use of ferrous sulphate and ferric chloride (the latter together with some glycerin). Iron salts have also been given with success in diphtheria, tonsilitis, and other forms of sore throat, as well as in erysipelas. In anæmia they have the effect of restoring the number of corpuscles and the normal amount of hæmoglobin in the blood. The fact that this takes place appears very remarkable in the light of a great deal of evidence which tends to show that no iron is actually absorbed into the system. We have seen above that in the stomach all iron salts are transformed into ferric chloride. On reaching the intestine the chloride is transformed into ferric hydroxide, and subsequently the latter is in turn transformed into the black sulphide and tannate of iron, which are voided with the fæces. All of the iron taken is thus voided, and none passes into the urine. On the other hand, when injected into the blood, even in very moderate quantities, iron salts produce symptoms of poisoning. The question therefore arises: In what manner do iron salts act in relieving amæmia? Definitely this question has not yet been answered. According to a theory advanced by Bunge, the iron normally present in the blood enters it in the form of complex organic iron compounds that are contained in food. That iron in some form or other necessarily enters the blood is evident, if we remember that the amount of iron in the body of a child increases with age. Now, according to Bunge, the alkaline sulphides that may be present in the intestines are capable of depriving the iron compounds of food of their iron, the resulting sulphide being of course incapable of absorption. But if sufficient quantities of iron are taken internally, the alkaline sulphides are decomposed and the organic iron of the food becomes available. The amounts of iron required depend of course upon the amount of alkaline sulphides in the intestines, and this is why it may be found necessary to administer as much as 18 grains a day to an anæmic woman whose body, in a normal state, contains altogether about 30 grains. A strong argument in favor of Bunge's theory of the indirect action of iron is found in the fact that manganese, copper, and certain other substances not at all present in the blood are almost as efficient as iron in curing anæmia.
Ferric chloride, the most important medicinal salt of iron, is usually administered in the form of its tincture, which contains about 3.25 per cent. of iron, corresponding to about 9.5 per cent. of anhydrous ferric chloride. The tincture is prepared by making up 250 parts of the official aqueous solution of ferric chloride to 1000 parts with alcohol. Ferric hydroxide with magnesia is known as ‘arsenic antidote,’ being an effective remedy for poisoning with arsenic. The antidote may be best prepared by gradually adding 10 parts of magnesia in water to 50 parts of ferric hydroxide in water, and shaking the mixture vigorously. It should be prepared immediately before using, and should be given repeatedly in large doses. Iron salts should never be given together with any preparation containing tannic or gallic acid.
IRON ORES.
The minerals which are commercially important as sources of iron may be grouped in the following classes: (1) Magnetite.—This class includes ores in which the iron occurs as magnetic oxide (Fe3O4); they contain when pure 72.40 per cent. of iron. (2) Hematite.—All varieties that are sesquioxides (Fe2O3), with 70 per cent. of iron. They are variously known as red, blue, and specular hematites, also as micaceous and fossil ores, according to their color and physical structure. (3) Limonite.—The hydrated oxides (2Fe2O33H2O), including bog ores, pipe ores, etc., with 59.89 per cent. of metal. Brown hematite is a synonymous term for limonite. (4) Siderite.—Ores containing carbon dioxide and represented by the type formula (FeCO3), with 48.27 per cent. of iron. Spathic ore is another name for siderite, while clay ironstone is a term applied to the varieties containing much clay and having a concretionary structure. When the ore contains bituminous matter in addition to clay it is called blackband.
Pyrites (FeS2) and franklinite, an oxide of iron, manganese, and zinc, are utilized to a very small extent. Pyrites is first roasted for the recovery of sulphur in sulphuric acid manufacture, and the clinker is then smelted. Franklinite is employed in the production of spiegeleisen, after extracting the zinc by roasting.
Composition. The ores of iron always contain more or less foreign matter, and only approximate the metallic content required by the chemical formulæ; the discrepancy commonly amounts to 10 per cent. or more. As the costs of handling and treatment per ton of iron are indirectly proportional to the purity of the ores, the higher grades are naturally in most demand. In the United States mining is confined practically to the hematite, limonite, and magnetite deposits, which on the average carry from 50 to 60 per cent. of iron, while elsewhere ores may be worked that run as low as 30 per cent. The nature of the impurities is of great importance in determining the value of ores. The common impurities are those which enter largely into the composition of the rocks surrounding the deposits. They are silica (SiO2), alumina (Al2O3), lime (CaO), magnesia (MgO), water (H2O), and carbon dioxide (CO2). Water and carbon dioxide are objectionable only as they replace the iron, while the others exert an influence upon the fluxing properties of the ore and the course of smelting operations. Of greater importance, however, are the small quantities of phosphorus, sulphur, and titanium, these impurities being almost wholly obnoxious in their effect. Phosphorus gives a fluid pig iron that can be converted into steel only by employing special methods of treatment. As a large part of the iron produced is now converted into steel by the acid Bessemer process, which does not eliminate phosphorus, ores adapted to this treatment find a ready market at good prices. For such ores the outside limit of phosphorus relative to the iron is 11000; that is, an ore carrying 60 per cent. of iron is not of Bessemer grade unless the phosphorous content falls below 0.06 per cent. In the United States common practice fixes a still lower limit for phosphorus.
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| COPYRIGHT, 1902, DODD, MEAD & COMPANY. |
Resources of the United States. The relative importance of the different iron ores mined in the United States is shown by the following percentages of the total output for 1901: Hematite, 83.1; limonite, 10.4; magnetite, 6.3; siderite, 0.2 per cent. Hematite thus contributes more than four-fifths of the total production. A large part of this ore comes from the Lake Superior region, where immense deposits have been found in metamorphosed Pre-Cambrian rocks. Five productive belts, or ‘ranges,’ as they are commonly called, are known. The Marquette range, opened in 1856, is situated in Michigan, east of the Keweenaw Peninsula. The Menominee Range, first developed in 1877, lies on the border of Wisconsin and the upper peninsula of Michigan; it is succeeded farther westward by the Gogebic Range, which was opened in 1884. The Vermilion and Mesabi ranges are situated northwest of Lake Superior, in Minnesota; they were first exploited in 1884. The deposits are found near the surface, and operations are conducted on a large scale. In some cases the ore is excavated, after removing the overburden, by steam shovels, which deposit it directly in the cars. A single mine in Minnesota has produced 1,681,000 long tons of ore in a year, or more than the annual output of the entire country prior to 1854. The total production of the Lake Superior mines up to and including 1901 was 192,008,000 long tons. The bulk of the ore is forwarded by rail to ports on Lakes Superior and Michigan, and thence shipped by boat to Chicago and Lake Erie ports, a large proportion being destined for the manufacturing centres of western Pennsylvania. A great shipping industry is engaged in the transport of Lake Superior ores, and it is largely owing to this fact that the tonnage of vessels passing through the Sault Sainte Marie Canal each year exceeds the annual commerce of the Suez Canal.
Hematite ores are also mined in many of the Appalachian and Rocky Mountain States. The Clinton formation of the Silurian, which outcrops along the western slopes of the Appalachians from New York to Alabama, contains valuable deposits interstratified with shales and limestones. Some ore is obtained at Clinton, N. Y., and near Chattanooga, Tenn.; but the most productive beds are in the Birmingham district of Alabama.
The limonite or bog ores are widely distributed, although, owing to their low iron content, they are worked extensively in only a few regions. Virginia, Alabama, Tennessee, and Colorado produce the largest quantities of limonite. Magnetite occurs in the Adirondacks, the New Jersey Highlands, near Cornwall, Pa., and in many of the Western States. Siderite is associated with the coal measures of Ohio and Pennsylvania.
The total production of iron ores in the United States in 1901, according to the Mineral Resources, was 28,887,479 long tons. This output was distributed among the leading States as follows: Minnesota, 11,109,537 tons; Michigan, 9,654,037 tons; Alabama, 2,801,732 tons; and Pennsylvania, 1,040,684 tons. The imports for the same year were 966,950 tons, more than one-half of which came from Cuba.
Foreign Countries. The iron ores of Great Britain include siderite, limonite, and hematite. The Cleveland district is the most important, and produces siderite averaging 30 to 40 per cent. in iron from deposits in the coal measures. Limonite is mined in Lincolnshire, Northamptonshire, and Leicestershire, and hematite in Lancashire and Cumberland. As the domestic supply of ore is insufficient, large quantities are imported from Spain, Sweden, Greece, and other countries. Germany has iron-mines in Alsace-Lorraine, Westphalia, and Hesse-Nassau, and the Grand Duchy of Luxemburg is noted for its iron. The ores are mostly limonites and hematites, and carry from 30 to 50 per cent. of iron. France, Spain, Sweden, Austria-Hungary, and Russia complete the list of the important producers. Canada has given much attention in recent years to the development of her iron industry, with results that promise well for the future. Large blast-furnaces and steel-works were completed in 1901, which utilize hematite and magnetite ores from Newfoundland, Nova Scotia, and Ontario.