cess, 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
