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, 55.9), 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 welded to gether 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, Fe2O; and ferric anhydride, FeO, or Fe2O. When oxidized at high temperatures, iron yields the substance Fe,O,; 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 Fe 0, which exhibits marked magnetic properties. The hydrate of ferric oxide, as ordinarily obtained by adding alkalies to solutions of ferric salts, has the composition 2Fe,0,.3H0, and is readily soluble in acids. Another hydrate of ferric oxide has the composition FeO.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 be 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, H.FeO,, which is likewise unknown in the free state, but certain of whose salts may be readily prepared. Thus, potassium ferrate, K.FeO,, 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 byproduct 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 FeSO,7H2O. 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 ox alate, 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 phos. phate. Ferric chloride, FeCl, 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), 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 Fe, (NO3) ̧.9H2O, and melt at 35° C. In aqueous solution, the nitrate gradually decomposes unless an excess of free nitric acid is present. Ferric phosphate, FePO,, 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 undesir able 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 anæmia? 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. 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 intes tines, 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.

That

Ferric chloride, the most important medicinal salt of iron, is usually administered in the form of its tincture, which contains about 3.25 per