Halogen Compounds.—Ferrous fluoride, FeF2, is obtained as colourless prisms (with 8H2O) by dissolving iron in hydrofluoric acid, or as anhydrous colourless rhombic prisms by heating iron or ferric chloride in dry hydrofluoric acid gas. Ferric fluoride, FeF3, is obtained as colourless crystals (with 4½H2O) by evaporating a solution of the hydroxide in hydrofluoric acid. When heated in air it yields ferric oxide. Ferrous chloride, FeCl2, is obtained as shining scales by passing chlorine, or, better, hydrochloric acid gas, over red-hot iron, or by reducing ferric chloride in a current of hydrogen. It is very deliquescent, and freely dissolves in water and alcohol. Heated in air it yields a mixture of ferric oxide and chloride, and in steam magnetic oxide, hydrochloric acid, and hydrogen. It absorbs ammonia gas, forming the compound FeCl2·6NH2, which on heating loses ammonia, and, finally, yields ammonium chloride, nitrogen and iron nitride. It fuses at a red-heat, and volatilizes at a yellow-heat; its vapour density at 1300°-1400° corresponds to the formula FeCl2. By evaporating in vacuo the solution obtained by dissolving iron in hydrochloric acid, there results bluish, monoclinic crystals of FeCl2·4H2O, which deliquesce, turning greenish, on exposure to air, and effloresce in a desiccator. Other hydrates are known. By adding ammonium chloride to the solution, evaporating in vacuo, and then volatilizing the ammonium chloride, anhydrous ferrous chloride is obtained. The solution, in common with those of most ferrous salts, absorbs nitric oxide with the formation of a brownish solution.Ferric chloride, FeCl3, known in its aqueous solution to Glauber asoleum martis, may be obtained anhydrous by the action of dry chlorine on the metal at a moderate red-heat, or by passing hydrochloric acid gas over heated ferric oxide. It forms iron-black plates or tablets which appear red by transmitted and a metallic green by reflected light. It is very deliquescent, and readily dissolves in water, forming a brown or yellow solution, from which several hydrates may be separated (seeSolution). The solution is best prepared by dissolving the hydrate in hydrochloric acid and removing the excess of acid by evaporation, or by passing chlorine into the solution obtained by dissolving the metal in hydrochloric acid and removing the excess of chlorine by a current of carbon dioxide. It also dissolves in alcohol and ether; boiling point determinations of the molecular weight in these solutions point to the formula FeCl3. Vapour density determinations at 448° indicate a partial dissociation of the double molecule Fe2Cl6; on stronger heating it splits into ferrous chloride and chlorine. It forms red crystalline double salts with the chlorides of the metals of the alkalis and of the magnesium group. An aqueous solution of ferric chloride is used in pharmacy under the nameLiquor ferri perchloridi; and an alcoholic solution constitutes the quack medicine known as “Lamotte’s golden drops.” Many oxychlorides are known; soluble forms are obtained by dissolving precipitated ferric hydrate in ferric chloride, whilst insoluble compounds result when ferrous chloride is oxidized in air, or by boiling for some time aqueous solutions of ferric chloride.Ferrous bromide, FeBr2, is obtained as yellowish crystals by the union of bromine and iron at a dull red-heat, or as bluish-green rhombic tables of the composition FeBr2·6H2O by crystallizing a solution of iron in hydrobromic acid. Ferric bromide, FeBr3, is obtained as dark red crystals by heating iron in an excess of bromine vapour. It closely resembles the chloride in being deliquescent, dissolving ferric hydrate, and in yielding basic salts. Ferrous iodide, FeI2, is obtained as a grey crystalline mass by the direct union of its components. Ferric iodide does not appear to exist.Sulphur Compounds.—Ferrous sulphide, FeS, results from the direct union of its elements, best by stirring molten sulphur with a white-hot iron rod, when the sulphide drops to the bottom of the crucible. It then forms a yellowish crystalline mass, which readily dissolves in acids with the liberation of sulphuretted hydrogen. Heated in air it at first partially oxidizes to ferrous sulphate, and at higher temperatures it yields sulphur dioxide and ferric oxide. It is unaltered by ignition in hydrogen. An amorphous form results when a mixture of iron filings and sulphur are triturated with water. This modification is rapidly oxidized by the air with such an elevation of temperature that the mass may become incandescent. Another black amorphous form results when ferrous salts are precipitated by ammonium sulphide.Ferric sulphide, Fe2S3, is obtained by gently heating a mixture of its constituent elements, or by the action of sulphuretted hydrogen on ferric oxide at temperatures below 100°. It is also prepared by precipitating a ferric salt with ammonium sulphide; unless the alkali be in excess a mixture of ferrous sulphide and sulphur is obtained. It combines with other sulphides to form compounds of the type M′2Fe2S4. Potassium ferric sulphide, K2Fe2S4, obtained by heating a mixture of iron filings, sulphur and potassium carbonate, forms purple glistening crystals, which burn when heated in air. Magnetic pyrites or pyrrhotite has a composition varying between Fe7S8and Fe8S9,i.e.5FeS·Fe2S3and 6FeS·Fe2S3. It has a somewhat brassy colour, and occurs massive or as hexagonal plates; it is attracted by a magnet and is sometimes itself magnetic. The mineral is abundant in Canada, where the presence of about 5% of nickel makes it a valuable ore of this metal. Iron disulphide, FeS2, constitutes the minerals pyrite and marcasite (q.v.); copper pyrites is (Cu, Fe)S2. Pyrite may be prepared artificially by gently heating ferrous sulphide with sulphur, or as brassy octahedra and cubes by slowly heating an intimate mixture of ferric oxide, sulphur and sal-ammoniac. It is insoluble in dilute acids, but dissolves in nitric acid with separation of sulphur.Ferrous sulphite, FeSO3. Iron dissolves in a solution of sulphur dioxide in the absence of air to form ferrous sulphite and thio-sulphate; the former, being less soluble than the latter, separates out as colourless or greenish crystals on standing.Ferrous sulphate, green vitriol or copperas, FeSO4·7H2O, was known to, and used by, the alchemists; it is mentioned in the writings of Agricola, and its preparation from iron and sulphuric acid occurs in theTractatus chymico-philosophicusascribed to Basil Valentine. It occurs in nature as the mineral melanterite, either crystalline or fibrous, but usually massive; it appears to have been formed by the oxidation of pyrite or marcasite. It is manufactured by piling pyrites in heaps and exposing to atmospheric oxidation, the ferrous sulphate thus formed being dissolved in water, and the solution run into tanks, where any sulphuric acid which may be formed is decomposed by adding scrap iron. By evaporation the green vitriol is obtained as large crystals. The chief impurities are copper and ferric sulphates; the former may be removed by adding scrap iron, which precipitates the copper; the latter is eliminated by recrystallization. Other impurities such as zinc and manganese sulphates are more difficult to remove, and hence to prepare the pure salt it is best to dissolve pure iron wire in dilute sulphuric acid. Ferrous sulphate forms large green crystals belonging to the monoclinic system; rhombic crystals, isomorphous with zinc sulphate, are obtained by inoculating a solution with a crystal of zinc sulphate, and triclinic crystals of the formula FeSO4·5H2O by inoculating with copper sulphate. By evaporating a solution containing free sulphuric acid in a vacuum, the hepta-hydrated salt first separates, then the penta-, and then a tetra-hydrate, FeSO4·4H2O, isomorphous with manganese sulphate. By gently heating in a vacuum to 140°, the hepta-hydrate loses 6 molecules of water, and yields a white powder, which on heating in the absence of air gives the anhydrous salt. The monohydrate also results as a white precipitate when concentrated sulphuric acid is added to a saturated solution of ferrous sulphate. Alcohol also throws down the salt from aqueous solution, the composition of the precipitate varying with the amount of salt and precipitant employed. The solution absorbs nitric oxide to form a dark brown solution, which loses the gas on heating or by placing in ä vacuum. Ferrous sulphate forms double salts with the alkaline sulphates. The most important is ferrous ammonium sulphate, FeSO4·(NH4)2SO4·6H2O, obtained by dissolving equivalent amountsof the two salts in water and crystallizing. It is very stable and is much used in volumetric analysis.Ferric sulphate, Fe2(SO4)3, is obtained by adding nitric acid to a hot solution of ferrous sulphate containing sulphuric acid, colourless crystals being deposited on evaporating the solution. The anhydrous salt is obtained by heating, or by adding concentrated sulphuric acid to a solution. It is sparingly soluble in water, and on heating it yields ferric oxide and sulphur dioxide. The mineral coquimbite is Fe2(SO4)3·9H2O. Many basic ferric sulphates are known, some of which occur as minerals; carphosiderite is Fe(FeO)5(SO4)4·10H2O; amarantite is Fe(FeO)(SO4)2·7H2O; utahite is 3(FeO)2SO4·4H2O; copiapite is Fe3(FeO)(SO4)5·18H2O; castanite is Fe(FeO)(SO4)2·8H2O; römerite is FeSO4·Fe2(SO4)3·12H2O. The iron alums are obtained by crystallizing solutions of equivalent quantities of ferric and an alkaline sulphate. Ferric potassium sulphate, the common iron alum, K2SO4·Fe2(SO4)3·24H2O, forms bright violet octahedra.Nitrides, Nitrates, &c.—Several nitrides are known. Guntz (Compt. rend., 1902, 135, p. 738) obtained ferrous nitride, Fe3N2, and ferric nitride, FeN, as black powders by heating lithium nitride with ferrous potassium chloride and ferric potassium chloride respectively. Fowler (Jour. Chem. Soc., 1901, p. 285) obtained a nitride Fe2N by acting upon anhydrous ferrous chloride or bromide, finely divided reduced iron, or iron amalgam with ammonia at 420°; and, also, in a compact form, by the action of ammonia on red-hot iron wire. It oxidizes on heating in air, and ignites in chlorine; on solution in mineral acids it yields ferrous and ammonium salts, hydrogen being liberated. A nitride appears to be formed when nitrogen is passed over heated iron, since the metal is rendered brittle. Ferrous nitrate, Fe(NO3)2·6H2O, is a very unstable salt, and is obtained by mixing solutions of ferrous sulphate and barium nitrate, filtering, and crystallizing in a vacuum over sulphuric acid. Ferric nitrate, Fe(NO3)3, is obtained by dissolving iron in nitric acid (the cold dilute acid leads to the formation of ferrous and ammonium nitrates) and crystallizing, when cubes of Fe(NO3)3·6H2O or monoclinic crystals of Fe(NO3)3·9H2O are obtained. It is used as a mordant.Ferrous solutions absorb nitric oxide, forming dark green to black solutions. The coloration is due to the production of unstable compounds of the ferrous salt and nitric oxide, and it seems that in neutral solutions the compound is made up of one molecule of salt to one of gas; the reaction, however, is reversible, the composition varying with temperature, concentration and nature of the salt. Ferrous chloride dissolved in strong hydrochloric acid absorbs two molecules of the gas (Kohlschütter and Kutscheroff,Ber., 1907, 40, p. 873). Ferric chloride also absorbs the gas. Reddish brown amorphous powders of the formulae 2FeCl3·NO and 4FeCl3·NO are obtained by passing the gas over anhydrous ferric chloride. By passing the gas into an ethereal solution of the salt, nitrosyl chloride is produced, and on evaporating over sulphuric acid, black needles of FeCl2·NO·2H2O are obtained, which at 60° form the yellow FeCl2·NO. Complicated compounds, discovered by Roussin in 1858, are obtained by the interaction of ferrous sulphate and alkaline nitrites and sulphides. Two classes may be distinguished:—(1) the ferrodinitroso salts,e.g.K[Fe(NO)2S], potassium ferrodinitrososulphide, and (2) the ferroheptanitroso salts,e.g.K[Fe4(NO)7S8], potassium ferroheptanitrososulphide. These salts yield the corresponding acids with sulphuric acid. The dinitroso acid slowly decomposes into sulphuretted hydrogen, nitrogen, nitrous oxide, and the heptanitroso acid. The heptanitroso acid is precipitated as a brown amorphous mass by dilute sulphuric acid, but if the salt be heated with strong acid it yields nitrogen, nitric oxide, sulphur, sulphuretted hydrogen, and ferric, ammonium and potassium sulphates.Phosphides, Phosphates.—H. Le Chatelier and S. Wologdine (Compt. rend., 1909, 149, p. 709) have obtained Fe3P, Fe2P, FeP, Fe2P3, but failed to prepare five other phosphides previously described. Fe3P occurs as crystals in the product of fusing iron with phosphorus; it dissolves in strong hydrochloric acid. Fe2P forms crystalline needles insoluble in acids except aqua regia; it is obtained by fusing copper phosphide with iron. FeP is obtained by passing phosphorus vapour over Fe2P at a red-heat. Fe2P3is prepared by the action of phosphorus iodide vapour on reduced iron. Ferrous phosphate, Fe3(PO4)2·8H2O, occurs in nature as the mineral vivianite. It may be obtained artificially as a white precipitate, which rapidly turns blue or green on exposure, by mixing solutions of ferrous sulphate and sodium phosphate. It is employed in medicine. Normal ferric phosphate, FePO4·2H2O, occurs as the mineral strengite, and is obtained as a yellowish-white precipitate by mixing solutions of ferric chloride and sodium phosphate. It is insoluble in dilute acetic acid, but dissolves in mineral acids. The acid salts Fe(H2PO4)3and 2FeH3(PO4)2·5H2O have been described. Basic salts have been prepared, and several occur in the mineral kingdom; dufrenite is Fe2(OH)3PO4.Arsenides, Arsenites, &c.—Several iron arsenides occur as minerals; lölingite, FeAs2, forms silvery rhombic prisms; mispickel or arsenical pyrites, Fe2AsS2, is an important commercial source of arsenic. A basic ferric arsenite, 4Fe2O3·As2O3·5H2O, is obtained as a flocculent brown precipitate by adding an arsenite to ferric acetate, or by shaking freshly prepared ferric hydrate with a solution of arsenious oxide. The last reaction is the basis of the application of ferric hydrate as an antidote in arsenical poisoning. Normal ferric arsenate, FeAsO4·2H2O, constitutes the mineral scorodite; pharmacosiderite is the basic arsenate 2FeAsO4·Fe(OK)3·5H2O. An acid arsenate, 2Fe2(HAsO4)3·9H2O, is obtained as a white precipitate by mixing solutions of ferric chloride and ordinary sodium phosphate. It readily dissolves in hydrochloric acid.Carbides, Carbonates.—The carbides of iron play an important part in determining the properties of the different modifications of the commercial metal, and are discussed underIron and Steel.Ferrous carbonate, FeCO3, or spathic iron ore, may be obtained as microscopic rhombohedra by adding sodium bicarbonate to ferrous sulphate and heating to 150° for 36 hours. Ferrous sulphate and sodium carbonate in the cold give a flocculent precipitate, at first white but rapidly turning green owing to oxidation. A soluble carbonate and a ferric salt give a precipitate which loses carbon dioxide on drying. Of great interest are the carbonyl compounds. Ferropentacarbonyl, Fe(CO)5, obtained by L. Mond, Quincke and Langer (Jour. Chem. Soc., 1891; see also ibid. 1910, p. 798) by treating iron from ferrous oxalate with carbon monoxide, and heating at 150°, is a pale yellow liquid which freezes at about −20°, and boils at 102.5°. Air and moisture decompose it. The halogens give ferrous and ferric haloids and carbon monoxide; hydrochloric and hydrobromic acids have no action, but hydriodic decomposes it. By exposure to sunlight, either alone or dissolved in ether or ligroin, it gives lustrous orange plates of diferrononacarbonyl, Fe2(CO)9. If this substance be heated in ethereal solution to 50°, it deposits lustrous dark-green tablets of ferrotetracarbonyl, Fe(CO)4, very stable at ordinary temperatures, but decomposing at 140°-150° into iron and carbon monoxide (J. Dewar and H. O. Jones,Abst. J.C.S., 1907, ii. 266). For the cyanides seePrussic Acid.Ferrous salts give a greenish precipitate with an alkali, whilst ferric give a characteristic red one. Ferrous salts also give a bluish white precipitate with ferrocyanide, which on exposure turns to a dark blue; ferric salts are characterized by the intense purple coloration with a thiocyanate. (See alsoChemistry, §Analytical). For the quantitative estimation seeAssaying.A recent atomic weight determination by Richards and Baxter (Zeit. anorg. Chem., 1900, 23, p. 245; 1904, 38, p. 232), who found the amount of silver bromide given by ferrous bromide, gave the value 55.44 [O = 16].Pharmacology.All the official salts and preparations of iron are made directly or indirectly from the metal. The pharmacopoeial forms of iron are as follow:—1.Ferrum, annealed iron wire No. 35 or wrought iron nails free from oxide; from which we have the preparationVinum ferri, iron wine, iron digested in sherry wine for thirty days. (Strength, 1 in 20.)2.Ferrum redactum, reduced iron, a powder containing at least 75% of metallic iron and a variable amount of oxide. A preparation of it isTrochiscus ferri redacti(strength, 1 grain of reduced iron in each).3.Ferri sulphas, ferrous sulphate, from which is preparedMistura ferri composita, “Griffiths’ mixture,” containing ferrous sulphate 25 gr., potassium carbonate 30 gr., myrrh 60 gr., sugar 60 gr., spirit of nutmeg 50 m., rose water 10 fl. oz.4.Ferri sulphas exsiccatus, which has two subpreparations: (a)Pilula ferri, “Blaud’s pill” (exsiccated ferrous sulphate 150, exsiccated sodium carbonate 95, gum acacia 50, tragacanth 15, glycerin 10, syrup 150, water 20, each to contain about 1 grain of ferrous carbonate); (b)Pilula aloes et ferri(Barbadoes aloes 2, exsiccated ferrous sulphate 1, compound powder of cinnamon 3, syrup of glucose 3).5.Ferri carbonas saccharatus, saccharated iron carbonate. The carbonate forms about one-third and is mixed with sugar into a greyish powder.6.Ferri arsenas, iron arsenate, ferrous and ferric arsenates with some iron oxides, a greenish powder.7.Ferri phosphas, a slate-blue powder of ferrous and ferric phosphates with some oxide. Its preparations are: (a)Syrupus ferri phosphatis(strength, 1 gr. of ferrous phosphate in each fluid drachm); (b)Syrupus ferri phosphatis cum quinina et strychnina, “Easton’s syrup” (iron wire 75 grs., concentrated phosphoric acid 10 fl. dr., powdered strychnine 5 gr., quinine sulphate 130 gr., syrup 14 fl. oz., water to make 20 fl. oz.), in which each fluid drachm represents 1 gr. of ferrous phosphate,4⁄5gr. of quinine sulphate, and1⁄32gr. of strychnine.8.Syrupus ferri iodidi, iron wire, iodine, water and syrup (strength, 5.5 gr. of ferrous iodide in one fl. dr.).9.Liquor ferri perchloridi fortis, strong solution of ferric chloride (strength, 22.5% of iron); its preparations only are prescribed, viz.Liquor ferri perchloridiandTinctura ferri perchloridi.10.Liquor ferri persulphatis, solution of ferric sulphate.11.Liquor ferri pernitratus, solution of ferric nitrate (strength, 3.3% of iron).12.Liquor ferri acetatis, solution of ferric acetate.13. The scale preparations of iron, so called because they are dried to form scales, are three in number, the base of all being ferric hydrate:(a)Ferrum tartaratum, dark red scales, soluble in water.(b)Ferri et quininae citratis, greenish yellow scales soluble in water.(c)Ferri et ammonii citratis, red scales soluble in water, from which is preparedVinum ferri citratis(ferri et ammonii citratis 1 gr., orange wine 1 fl. dr.).Substances containing tannic or gallic acid turn black when compounded with a ferric salt, so it cannot be used in combination with vegetable astringents except with the infusion of quassia or calumba. Iron may, however, be prescribed in combination with digitalis by the addition of dilute phosphoric acid. Alkalis and their carbonates, lime water, carbonate of calcium, magnesia and its carbonate give green precipitates with ferrous and brown with ferric salts.Unofficial preparations of iron are numberless, and some of them are very useful. Ferri hydroxidum (U.S.P.), the hydrated oxide of iron, made by precipitating ferric sulphate with ammonia, is used solely as an antidote in arsenical poisoning. The Syrupus ferri phosphatis Co. is well known as “Parrish’s” syrup or chemical food, and the Pilulae ferri phosphatis cum quinina et strychnina, known as Easton’s pills, form a solid equivalent to Easton’s syrup.There are numerous organic preparations of iron. Ferratin is a reddish brown substance which claims to be identical with the iron substance found in pig’s liver. Carniferrin is another tasteless powder containing iron in combination with the phosphocarnic acid of muscle preparations, and contains 35% of iron. Ferratogen is prepared from ferric nuclein. Triferrin is a paranucleinate of iron, and contains 22% of iron and 2½% of organically combined phosphorus, prepared from the casein of cow’s milk. Haemoglobin is extracted from the blood of an ox and may be administered in bolus form. Dieterich’s solution of peptonated iron contains about 2 gr. of iron per oz. Vachetta has used the albuminate of iron with striking success in grave cases of anaemia. Succinate of iron has been prepared by Hausmann. Haematogen, introduced by Hommel, claims to contain the albuminous constituents of the blood serum and all the blood salts as well as pure haemoglobin. Sicco, the name given to dry haematogen, is a tasteless powder. Haemalbumen, introduced by Dahmen, is soluble in warm water.Therapeutics.Iron is a metal which is used both as a food and as a medicine and has also a definite local action. Externally, it is not absorbed by the unbroken skin, but when applied to the broken skin, sores, ulcers and mucous surfaces, the ferric salts are powerful astringents, because they coagulate the albuminous fluids in the tissues themselves. The salts of iron quickly cause coagulation of the blood, and the clot plugs the bleeding vessels. They thus act locally as haemostatics or styptics, and will often arrest severe haemorrhage from parts which are accessible, such as the nose. They were formerly used in the treatment ofpost partumhaemorrhage. The perchloride, sulphate and pernitrate are strongly astringent; less extensively they are used in chronic discharges from the vagina, rectum and nose, while injected into the rectum they destroy worms.Internally, a large proportion of the various articles of ordinary diet contains iron. When given medicinally preparations of iron have an astringent taste, and the teeth and tongue are blackened owing to the formation of sulphide of iron. It is therefore advisable to take liquid iron preparations through a glass tube or a quill.In the stomach all salts of iron, whatever their nature, are converted into ferric chloride. If iron be given in excess, or if the hydrochloric acid in the gastric juice be deficient, iron acts directly as an astringent upon the mucous membrane of the stomach wall. Iron, therefore, may disorder the digestion even in healthy subjects. Acid preparations are more likely to do this, and the acid set free after the formation of the chloride may act as an irritant. Iron, therefore, must not be given to subjects in whom the gastric functions are disturbed, and it should always be given after meals. Preparations which are not acid, or are only slightly acid, such as reduced iron, dialysed iron, the carbonate and scale preparations, do not disturb the digestion. If the sulphate is prescribed in the form of a pill, it may be so coated as only to be soluble in the intestinal digestive fluid. In the intestine the ferric chloride becomes changed into an oxide of iron; the sub-chloride is converted into a ferrous carbonate, which is soluble. Lower down in the bowel these compounds are converted into ferrous sulphide and tannate, and are eliminated with the faeces, turning them black. Iron in the intestine causes an astringent or constipating effect. The astringent salts are therefore useful occasionally to check diarrhoea and dysentery. Thus most salts of iron are distinctly constipating, and are best used in combination with a purgative. The pill of iron and aloes (B.P.) is designed for this purpose. Iron is certainly absorbed from the intestinal canal. As the iron in the food supplies all the iron in the body of a healthy person, there is no doubt that it is absorbed in the organic form. Whether inorganic salts are directly absorbed has been a matter of much discussion; it has, however, been directly proved by the experiments of Kunkel (Archiv für die gesamte Physiologie des Menschen und der Tiere, lxi.) and Gaule. The amount of iron existing in the human blood is only 38 gr.; therefore, when an excess of iron is absorbed, part is excreted immediately by the bowel and kidneys, and part is stored in the liver and spleen.Iron being a constituent part of the blood itself, there is a direct indication for the physician to prescribe it when the amount of haemoglobin in the blood is lowered or the red corpuscles are diminished. In certain forms of anaemia the administration of iron rapidly improves the blood in both respects. The exact method in which the prescribed iron acts is still a matter of dispute. Ralph Stockman points out that there are three chief theories as to the action of iron in anaemia. The first is based on the fact that the iron in the haemoglobin of the blood must be derived from the food, therefore iron medicinally administered is absorbed. The second theory is that there is no absorption of iron given by the mouth, but it acts as a local stimulant to the mucous membrane, and so improves anaemia by increasing the digestion of the food. The third theory is that of Bunge, who says that in chlorotic conditions there is an excess of sulphuretted hydrogen in the bowel, changing the food iron into sulphide of iron, which Bunge states cannot be absorbed. He believes that inorganic iron saves the organic iron of the food by combining with the sulphur, and improves anaemia by protecting the organic food iron. Stockman’s own experiments are, however, directly opposed to Bunge’s view. Wharfinger states that in chlorosis the specific action of iron is only obtained by administering those inorganic preparations which give a reaction with the ordinary reagents; the iron ions in a state of dissociation act as a catalytic agent, destroying the hypothetical toxin which is the cause of chlorosis. Practical experience teaches every clinician that, whatever the mode of action, iron is most valuable in anaemia, though in many cases, where there is well-marked toxaemia from absorption of the intestinal products, not only laxatives in combination with iron but intestinal antiseptics are necessary. That form of neuralgia which is associated with anaemia usually yields to iron.
Halogen Compounds.—Ferrous fluoride, FeF2, is obtained as colourless prisms (with 8H2O) by dissolving iron in hydrofluoric acid, or as anhydrous colourless rhombic prisms by heating iron or ferric chloride in dry hydrofluoric acid gas. Ferric fluoride, FeF3, is obtained as colourless crystals (with 4½H2O) by evaporating a solution of the hydroxide in hydrofluoric acid. When heated in air it yields ferric oxide. Ferrous chloride, FeCl2, is obtained as shining scales by passing chlorine, or, better, hydrochloric acid gas, over red-hot iron, or by reducing ferric chloride in a current of hydrogen. It is very deliquescent, and freely dissolves in water and alcohol. Heated in air it yields a mixture of ferric oxide and chloride, and in steam magnetic oxide, hydrochloric acid, and hydrogen. It absorbs ammonia gas, forming the compound FeCl2·6NH2, which on heating loses ammonia, and, finally, yields ammonium chloride, nitrogen and iron nitride. It fuses at a red-heat, and volatilizes at a yellow-heat; its vapour density at 1300°-1400° corresponds to the formula FeCl2. By evaporating in vacuo the solution obtained by dissolving iron in hydrochloric acid, there results bluish, monoclinic crystals of FeCl2·4H2O, which deliquesce, turning greenish, on exposure to air, and effloresce in a desiccator. Other hydrates are known. By adding ammonium chloride to the solution, evaporating in vacuo, and then volatilizing the ammonium chloride, anhydrous ferrous chloride is obtained. The solution, in common with those of most ferrous salts, absorbs nitric oxide with the formation of a brownish solution.
Ferric chloride, FeCl3, known in its aqueous solution to Glauber asoleum martis, may be obtained anhydrous by the action of dry chlorine on the metal at a moderate red-heat, or by passing hydrochloric acid gas over heated ferric oxide. It forms iron-black plates or tablets which appear red by transmitted and a metallic green by reflected light. It is very deliquescent, and readily dissolves in water, forming a brown or yellow solution, from which several hydrates may be separated (seeSolution). The solution is best prepared by dissolving the hydrate in hydrochloric acid and removing the excess of acid by evaporation, or by passing chlorine into the solution obtained by dissolving the metal in hydrochloric acid and removing the excess of chlorine by a current of carbon dioxide. It also dissolves in alcohol and ether; boiling point determinations of the molecular weight in these solutions point to the formula FeCl3. Vapour density determinations at 448° indicate a partial dissociation of the double molecule Fe2Cl6; on stronger heating it splits into ferrous chloride and chlorine. It forms red crystalline double salts with the chlorides of the metals of the alkalis and of the magnesium group. An aqueous solution of ferric chloride is used in pharmacy under the nameLiquor ferri perchloridi; and an alcoholic solution constitutes the quack medicine known as “Lamotte’s golden drops.” Many oxychlorides are known; soluble forms are obtained by dissolving precipitated ferric hydrate in ferric chloride, whilst insoluble compounds result when ferrous chloride is oxidized in air, or by boiling for some time aqueous solutions of ferric chloride.
Ferrous bromide, FeBr2, is obtained as yellowish crystals by the union of bromine and iron at a dull red-heat, or as bluish-green rhombic tables of the composition FeBr2·6H2O by crystallizing a solution of iron in hydrobromic acid. Ferric bromide, FeBr3, is obtained as dark red crystals by heating iron in an excess of bromine vapour. It closely resembles the chloride in being deliquescent, dissolving ferric hydrate, and in yielding basic salts. Ferrous iodide, FeI2, is obtained as a grey crystalline mass by the direct union of its components. Ferric iodide does not appear to exist.
Sulphur Compounds.—Ferrous sulphide, FeS, results from the direct union of its elements, best by stirring molten sulphur with a white-hot iron rod, when the sulphide drops to the bottom of the crucible. It then forms a yellowish crystalline mass, which readily dissolves in acids with the liberation of sulphuretted hydrogen. Heated in air it at first partially oxidizes to ferrous sulphate, and at higher temperatures it yields sulphur dioxide and ferric oxide. It is unaltered by ignition in hydrogen. An amorphous form results when a mixture of iron filings and sulphur are triturated with water. This modification is rapidly oxidized by the air with such an elevation of temperature that the mass may become incandescent. Another black amorphous form results when ferrous salts are precipitated by ammonium sulphide.
Ferric sulphide, Fe2S3, is obtained by gently heating a mixture of its constituent elements, or by the action of sulphuretted hydrogen on ferric oxide at temperatures below 100°. It is also prepared by precipitating a ferric salt with ammonium sulphide; unless the alkali be in excess a mixture of ferrous sulphide and sulphur is obtained. It combines with other sulphides to form compounds of the type M′2Fe2S4. Potassium ferric sulphide, K2Fe2S4, obtained by heating a mixture of iron filings, sulphur and potassium carbonate, forms purple glistening crystals, which burn when heated in air. Magnetic pyrites or pyrrhotite has a composition varying between Fe7S8and Fe8S9,i.e.5FeS·Fe2S3and 6FeS·Fe2S3. It has a somewhat brassy colour, and occurs massive or as hexagonal plates; it is attracted by a magnet and is sometimes itself magnetic. The mineral is abundant in Canada, where the presence of about 5% of nickel makes it a valuable ore of this metal. Iron disulphide, FeS2, constitutes the minerals pyrite and marcasite (q.v.); copper pyrites is (Cu, Fe)S2. Pyrite may be prepared artificially by gently heating ferrous sulphide with sulphur, or as brassy octahedra and cubes by slowly heating an intimate mixture of ferric oxide, sulphur and sal-ammoniac. It is insoluble in dilute acids, but dissolves in nitric acid with separation of sulphur.
Ferrous sulphite, FeSO3. Iron dissolves in a solution of sulphur dioxide in the absence of air to form ferrous sulphite and thio-sulphate; the former, being less soluble than the latter, separates out as colourless or greenish crystals on standing.
Ferrous sulphate, green vitriol or copperas, FeSO4·7H2O, was known to, and used by, the alchemists; it is mentioned in the writings of Agricola, and its preparation from iron and sulphuric acid occurs in theTractatus chymico-philosophicusascribed to Basil Valentine. It occurs in nature as the mineral melanterite, either crystalline or fibrous, but usually massive; it appears to have been formed by the oxidation of pyrite or marcasite. It is manufactured by piling pyrites in heaps and exposing to atmospheric oxidation, the ferrous sulphate thus formed being dissolved in water, and the solution run into tanks, where any sulphuric acid which may be formed is decomposed by adding scrap iron. By evaporation the green vitriol is obtained as large crystals. The chief impurities are copper and ferric sulphates; the former may be removed by adding scrap iron, which precipitates the copper; the latter is eliminated by recrystallization. Other impurities such as zinc and manganese sulphates are more difficult to remove, and hence to prepare the pure salt it is best to dissolve pure iron wire in dilute sulphuric acid. Ferrous sulphate forms large green crystals belonging to the monoclinic system; rhombic crystals, isomorphous with zinc sulphate, are obtained by inoculating a solution with a crystal of zinc sulphate, and triclinic crystals of the formula FeSO4·5H2O by inoculating with copper sulphate. By evaporating a solution containing free sulphuric acid in a vacuum, the hepta-hydrated salt first separates, then the penta-, and then a tetra-hydrate, FeSO4·4H2O, isomorphous with manganese sulphate. By gently heating in a vacuum to 140°, the hepta-hydrate loses 6 molecules of water, and yields a white powder, which on heating in the absence of air gives the anhydrous salt. The monohydrate also results as a white precipitate when concentrated sulphuric acid is added to a saturated solution of ferrous sulphate. Alcohol also throws down the salt from aqueous solution, the composition of the precipitate varying with the amount of salt and precipitant employed. The solution absorbs nitric oxide to form a dark brown solution, which loses the gas on heating or by placing in ä vacuum. Ferrous sulphate forms double salts with the alkaline sulphates. The most important is ferrous ammonium sulphate, FeSO4·(NH4)2SO4·6H2O, obtained by dissolving equivalent amountsof the two salts in water and crystallizing. It is very stable and is much used in volumetric analysis.
Ferric sulphate, Fe2(SO4)3, is obtained by adding nitric acid to a hot solution of ferrous sulphate containing sulphuric acid, colourless crystals being deposited on evaporating the solution. The anhydrous salt is obtained by heating, or by adding concentrated sulphuric acid to a solution. It is sparingly soluble in water, and on heating it yields ferric oxide and sulphur dioxide. The mineral coquimbite is Fe2(SO4)3·9H2O. Many basic ferric sulphates are known, some of which occur as minerals; carphosiderite is Fe(FeO)5(SO4)4·10H2O; amarantite is Fe(FeO)(SO4)2·7H2O; utahite is 3(FeO)2SO4·4H2O; copiapite is Fe3(FeO)(SO4)5·18H2O; castanite is Fe(FeO)(SO4)2·8H2O; römerite is FeSO4·Fe2(SO4)3·12H2O. The iron alums are obtained by crystallizing solutions of equivalent quantities of ferric and an alkaline sulphate. Ferric potassium sulphate, the common iron alum, K2SO4·Fe2(SO4)3·24H2O, forms bright violet octahedra.
Nitrides, Nitrates, &c.—Several nitrides are known. Guntz (Compt. rend., 1902, 135, p. 738) obtained ferrous nitride, Fe3N2, and ferric nitride, FeN, as black powders by heating lithium nitride with ferrous potassium chloride and ferric potassium chloride respectively. Fowler (Jour. Chem. Soc., 1901, p. 285) obtained a nitride Fe2N by acting upon anhydrous ferrous chloride or bromide, finely divided reduced iron, or iron amalgam with ammonia at 420°; and, also, in a compact form, by the action of ammonia on red-hot iron wire. It oxidizes on heating in air, and ignites in chlorine; on solution in mineral acids it yields ferrous and ammonium salts, hydrogen being liberated. A nitride appears to be formed when nitrogen is passed over heated iron, since the metal is rendered brittle. Ferrous nitrate, Fe(NO3)2·6H2O, is a very unstable salt, and is obtained by mixing solutions of ferrous sulphate and barium nitrate, filtering, and crystallizing in a vacuum over sulphuric acid. Ferric nitrate, Fe(NO3)3, is obtained by dissolving iron in nitric acid (the cold dilute acid leads to the formation of ferrous and ammonium nitrates) and crystallizing, when cubes of Fe(NO3)3·6H2O or monoclinic crystals of Fe(NO3)3·9H2O are obtained. It is used as a mordant.
Ferrous solutions absorb nitric oxide, forming dark green to black solutions. The coloration is due to the production of unstable compounds of the ferrous salt and nitric oxide, and it seems that in neutral solutions the compound is made up of one molecule of salt to one of gas; the reaction, however, is reversible, the composition varying with temperature, concentration and nature of the salt. Ferrous chloride dissolved in strong hydrochloric acid absorbs two molecules of the gas (Kohlschütter and Kutscheroff,Ber., 1907, 40, p. 873). Ferric chloride also absorbs the gas. Reddish brown amorphous powders of the formulae 2FeCl3·NO and 4FeCl3·NO are obtained by passing the gas over anhydrous ferric chloride. By passing the gas into an ethereal solution of the salt, nitrosyl chloride is produced, and on evaporating over sulphuric acid, black needles of FeCl2·NO·2H2O are obtained, which at 60° form the yellow FeCl2·NO. Complicated compounds, discovered by Roussin in 1858, are obtained by the interaction of ferrous sulphate and alkaline nitrites and sulphides. Two classes may be distinguished:—(1) the ferrodinitroso salts,e.g.K[Fe(NO)2S], potassium ferrodinitrososulphide, and (2) the ferroheptanitroso salts,e.g.K[Fe4(NO)7S8], potassium ferroheptanitrososulphide. These salts yield the corresponding acids with sulphuric acid. The dinitroso acid slowly decomposes into sulphuretted hydrogen, nitrogen, nitrous oxide, and the heptanitroso acid. The heptanitroso acid is precipitated as a brown amorphous mass by dilute sulphuric acid, but if the salt be heated with strong acid it yields nitrogen, nitric oxide, sulphur, sulphuretted hydrogen, and ferric, ammonium and potassium sulphates.
Phosphides, Phosphates.—H. Le Chatelier and S. Wologdine (Compt. rend., 1909, 149, p. 709) have obtained Fe3P, Fe2P, FeP, Fe2P3, but failed to prepare five other phosphides previously described. Fe3P occurs as crystals in the product of fusing iron with phosphorus; it dissolves in strong hydrochloric acid. Fe2P forms crystalline needles insoluble in acids except aqua regia; it is obtained by fusing copper phosphide with iron. FeP is obtained by passing phosphorus vapour over Fe2P at a red-heat. Fe2P3is prepared by the action of phosphorus iodide vapour on reduced iron. Ferrous phosphate, Fe3(PO4)2·8H2O, occurs in nature as the mineral vivianite. It may be obtained artificially as a white precipitate, which rapidly turns blue or green on exposure, by mixing solutions of ferrous sulphate and sodium phosphate. It is employed in medicine. Normal ferric phosphate, FePO4·2H2O, occurs as the mineral strengite, and is obtained as a yellowish-white precipitate by mixing solutions of ferric chloride and sodium phosphate. It is insoluble in dilute acetic acid, but dissolves in mineral acids. The acid salts Fe(H2PO4)3and 2FeH3(PO4)2·5H2O have been described. Basic salts have been prepared, and several occur in the mineral kingdom; dufrenite is Fe2(OH)3PO4.
Arsenides, Arsenites, &c.—Several iron arsenides occur as minerals; lölingite, FeAs2, forms silvery rhombic prisms; mispickel or arsenical pyrites, Fe2AsS2, is an important commercial source of arsenic. A basic ferric arsenite, 4Fe2O3·As2O3·5H2O, is obtained as a flocculent brown precipitate by adding an arsenite to ferric acetate, or by shaking freshly prepared ferric hydrate with a solution of arsenious oxide. The last reaction is the basis of the application of ferric hydrate as an antidote in arsenical poisoning. Normal ferric arsenate, FeAsO4·2H2O, constitutes the mineral scorodite; pharmacosiderite is the basic arsenate 2FeAsO4·Fe(OK)3·5H2O. An acid arsenate, 2Fe2(HAsO4)3·9H2O, is obtained as a white precipitate by mixing solutions of ferric chloride and ordinary sodium phosphate. It readily dissolves in hydrochloric acid.
Carbides, Carbonates.—The carbides of iron play an important part in determining the properties of the different modifications of the commercial metal, and are discussed underIron and Steel.
Ferrous carbonate, FeCO3, or spathic iron ore, may be obtained as microscopic rhombohedra by adding sodium bicarbonate to ferrous sulphate and heating to 150° for 36 hours. Ferrous sulphate and sodium carbonate in the cold give a flocculent precipitate, at first white but rapidly turning green owing to oxidation. A soluble carbonate and a ferric salt give a precipitate which loses carbon dioxide on drying. Of great interest are the carbonyl compounds. Ferropentacarbonyl, Fe(CO)5, obtained by L. Mond, Quincke and Langer (Jour. Chem. Soc., 1891; see also ibid. 1910, p. 798) by treating iron from ferrous oxalate with carbon monoxide, and heating at 150°, is a pale yellow liquid which freezes at about −20°, and boils at 102.5°. Air and moisture decompose it. The halogens give ferrous and ferric haloids and carbon monoxide; hydrochloric and hydrobromic acids have no action, but hydriodic decomposes it. By exposure to sunlight, either alone or dissolved in ether or ligroin, it gives lustrous orange plates of diferrononacarbonyl, Fe2(CO)9. If this substance be heated in ethereal solution to 50°, it deposits lustrous dark-green tablets of ferrotetracarbonyl, Fe(CO)4, very stable at ordinary temperatures, but decomposing at 140°-150° into iron and carbon monoxide (J. Dewar and H. O. Jones,Abst. J.C.S., 1907, ii. 266). For the cyanides seePrussic Acid.
Ferrous salts give a greenish precipitate with an alkali, whilst ferric give a characteristic red one. Ferrous salts also give a bluish white precipitate with ferrocyanide, which on exposure turns to a dark blue; ferric salts are characterized by the intense purple coloration with a thiocyanate. (See alsoChemistry, §Analytical). For the quantitative estimation seeAssaying.
A recent atomic weight determination by Richards and Baxter (Zeit. anorg. Chem., 1900, 23, p. 245; 1904, 38, p. 232), who found the amount of silver bromide given by ferrous bromide, gave the value 55.44 [O = 16].
Pharmacology.
All the official salts and preparations of iron are made directly or indirectly from the metal. The pharmacopoeial forms of iron are as follow:—
1.Ferrum, annealed iron wire No. 35 or wrought iron nails free from oxide; from which we have the preparationVinum ferri, iron wine, iron digested in sherry wine for thirty days. (Strength, 1 in 20.)
2.Ferrum redactum, reduced iron, a powder containing at least 75% of metallic iron and a variable amount of oxide. A preparation of it isTrochiscus ferri redacti(strength, 1 grain of reduced iron in each).
3.Ferri sulphas, ferrous sulphate, from which is preparedMistura ferri composita, “Griffiths’ mixture,” containing ferrous sulphate 25 gr., potassium carbonate 30 gr., myrrh 60 gr., sugar 60 gr., spirit of nutmeg 50 m., rose water 10 fl. oz.
4.Ferri sulphas exsiccatus, which has two subpreparations: (a)Pilula ferri, “Blaud’s pill” (exsiccated ferrous sulphate 150, exsiccated sodium carbonate 95, gum acacia 50, tragacanth 15, glycerin 10, syrup 150, water 20, each to contain about 1 grain of ferrous carbonate); (b)Pilula aloes et ferri(Barbadoes aloes 2, exsiccated ferrous sulphate 1, compound powder of cinnamon 3, syrup of glucose 3).
5.Ferri carbonas saccharatus, saccharated iron carbonate. The carbonate forms about one-third and is mixed with sugar into a greyish powder.
6.Ferri arsenas, iron arsenate, ferrous and ferric arsenates with some iron oxides, a greenish powder.
7.Ferri phosphas, a slate-blue powder of ferrous and ferric phosphates with some oxide. Its preparations are: (a)Syrupus ferri phosphatis(strength, 1 gr. of ferrous phosphate in each fluid drachm); (b)Syrupus ferri phosphatis cum quinina et strychnina, “Easton’s syrup” (iron wire 75 grs., concentrated phosphoric acid 10 fl. dr., powdered strychnine 5 gr., quinine sulphate 130 gr., syrup 14 fl. oz., water to make 20 fl. oz.), in which each fluid drachm represents 1 gr. of ferrous phosphate,4⁄5gr. of quinine sulphate, and1⁄32gr. of strychnine.
8.Syrupus ferri iodidi, iron wire, iodine, water and syrup (strength, 5.5 gr. of ferrous iodide in one fl. dr.).
9.Liquor ferri perchloridi fortis, strong solution of ferric chloride (strength, 22.5% of iron); its preparations only are prescribed, viz.Liquor ferri perchloridiandTinctura ferri perchloridi.
10.Liquor ferri persulphatis, solution of ferric sulphate.
11.Liquor ferri pernitratus, solution of ferric nitrate (strength, 3.3% of iron).
12.Liquor ferri acetatis, solution of ferric acetate.
13. The scale preparations of iron, so called because they are dried to form scales, are three in number, the base of all being ferric hydrate:
(a)Ferrum tartaratum, dark red scales, soluble in water.
(b)Ferri et quininae citratis, greenish yellow scales soluble in water.
(c)Ferri et ammonii citratis, red scales soluble in water, from which is preparedVinum ferri citratis(ferri et ammonii citratis 1 gr., orange wine 1 fl. dr.).
Substances containing tannic or gallic acid turn black when compounded with a ferric salt, so it cannot be used in combination with vegetable astringents except with the infusion of quassia or calumba. Iron may, however, be prescribed in combination with digitalis by the addition of dilute phosphoric acid. Alkalis and their carbonates, lime water, carbonate of calcium, magnesia and its carbonate give green precipitates with ferrous and brown with ferric salts.
Unofficial preparations of iron are numberless, and some of them are very useful. Ferri hydroxidum (U.S.P.), the hydrated oxide of iron, made by precipitating ferric sulphate with ammonia, is used solely as an antidote in arsenical poisoning. The Syrupus ferri phosphatis Co. is well known as “Parrish’s” syrup or chemical food, and the Pilulae ferri phosphatis cum quinina et strychnina, known as Easton’s pills, form a solid equivalent to Easton’s syrup.
There are numerous organic preparations of iron. Ferratin is a reddish brown substance which claims to be identical with the iron substance found in pig’s liver. Carniferrin is another tasteless powder containing iron in combination with the phosphocarnic acid of muscle preparations, and contains 35% of iron. Ferratogen is prepared from ferric nuclein. Triferrin is a paranucleinate of iron, and contains 22% of iron and 2½% of organically combined phosphorus, prepared from the casein of cow’s milk. Haemoglobin is extracted from the blood of an ox and may be administered in bolus form. Dieterich’s solution of peptonated iron contains about 2 gr. of iron per oz. Vachetta has used the albuminate of iron with striking success in grave cases of anaemia. Succinate of iron has been prepared by Hausmann. Haematogen, introduced by Hommel, claims to contain the albuminous constituents of the blood serum and all the blood salts as well as pure haemoglobin. Sicco, the name given to dry haematogen, is a tasteless powder. Haemalbumen, introduced by Dahmen, is soluble in warm water.
Therapeutics.
Iron is a metal which is used both as a food and as a medicine and has also a definite local action. Externally, it is not absorbed by the unbroken skin, but when applied to the broken skin, sores, ulcers and mucous surfaces, the ferric salts are powerful astringents, because they coagulate the albuminous fluids in the tissues themselves. The salts of iron quickly cause coagulation of the blood, and the clot plugs the bleeding vessels. They thus act locally as haemostatics or styptics, and will often arrest severe haemorrhage from parts which are accessible, such as the nose. They were formerly used in the treatment ofpost partumhaemorrhage. The perchloride, sulphate and pernitrate are strongly astringent; less extensively they are used in chronic discharges from the vagina, rectum and nose, while injected into the rectum they destroy worms.
Internally, a large proportion of the various articles of ordinary diet contains iron. When given medicinally preparations of iron have an astringent taste, and the teeth and tongue are blackened owing to the formation of sulphide of iron. It is therefore advisable to take liquid iron preparations through a glass tube or a quill.
In the stomach all salts of iron, whatever their nature, are converted into ferric chloride. If iron be given in excess, or if the hydrochloric acid in the gastric juice be deficient, iron acts directly as an astringent upon the mucous membrane of the stomach wall. Iron, therefore, may disorder the digestion even in healthy subjects. Acid preparations are more likely to do this, and the acid set free after the formation of the chloride may act as an irritant. Iron, therefore, must not be given to subjects in whom the gastric functions are disturbed, and it should always be given after meals. Preparations which are not acid, or are only slightly acid, such as reduced iron, dialysed iron, the carbonate and scale preparations, do not disturb the digestion. If the sulphate is prescribed in the form of a pill, it may be so coated as only to be soluble in the intestinal digestive fluid. In the intestine the ferric chloride becomes changed into an oxide of iron; the sub-chloride is converted into a ferrous carbonate, which is soluble. Lower down in the bowel these compounds are converted into ferrous sulphide and tannate, and are eliminated with the faeces, turning them black. Iron in the intestine causes an astringent or constipating effect. The astringent salts are therefore useful occasionally to check diarrhoea and dysentery. Thus most salts of iron are distinctly constipating, and are best used in combination with a purgative. The pill of iron and aloes (B.P.) is designed for this purpose. Iron is certainly absorbed from the intestinal canal. As the iron in the food supplies all the iron in the body of a healthy person, there is no doubt that it is absorbed in the organic form. Whether inorganic salts are directly absorbed has been a matter of much discussion; it has, however, been directly proved by the experiments of Kunkel (Archiv für die gesamte Physiologie des Menschen und der Tiere, lxi.) and Gaule. The amount of iron existing in the human blood is only 38 gr.; therefore, when an excess of iron is absorbed, part is excreted immediately by the bowel and kidneys, and part is stored in the liver and spleen.
Iron being a constituent part of the blood itself, there is a direct indication for the physician to prescribe it when the amount of haemoglobin in the blood is lowered or the red corpuscles are diminished. In certain forms of anaemia the administration of iron rapidly improves the blood in both respects. The exact method in which the prescribed iron acts is still a matter of dispute. Ralph Stockman points out that there are three chief theories as to the action of iron in anaemia. The first is based on the fact that the iron in the haemoglobin of the blood must be derived from the food, therefore iron medicinally administered is absorbed. The second theory is that there is no absorption of iron given by the mouth, but it acts as a local stimulant to the mucous membrane, and so improves anaemia by increasing the digestion of the food. The third theory is that of Bunge, who says that in chlorotic conditions there is an excess of sulphuretted hydrogen in the bowel, changing the food iron into sulphide of iron, which Bunge states cannot be absorbed. He believes that inorganic iron saves the organic iron of the food by combining with the sulphur, and improves anaemia by protecting the organic food iron. Stockman’s own experiments are, however, directly opposed to Bunge’s view. Wharfinger states that in chlorosis the specific action of iron is only obtained by administering those inorganic preparations which give a reaction with the ordinary reagents; the iron ions in a state of dissociation act as a catalytic agent, destroying the hypothetical toxin which is the cause of chlorosis. Practical experience teaches every clinician that, whatever the mode of action, iron is most valuable in anaemia, though in many cases, where there is well-marked toxaemia from absorption of the intestinal products, not only laxatives in combination with iron but intestinal antiseptics are necessary. That form of neuralgia which is associated with anaemia usually yields to iron.
1By solution in concentrated hydrochloric acid, a yellow liquid is obtained, which on concentration over sulphuric acid gives yellow deliquescent crusts of ferroso-ferric chloride, Fe3Cl8·18H2O.
1By solution in concentrated hydrochloric acid, a yellow liquid is obtained, which on concentration over sulphuric acid gives yellow deliquescent crusts of ferroso-ferric chloride, Fe3Cl8·18H2O.
IRON AGE,the third of the three periods, Stone, Bronze and Iron Ages, into which archaeologists divide prehistoric time; the weapons, utensils and implements being as a general rule made of iron (seeArchaeology). The term has no real chronological value, for there has been no universal synchronous sequence of the three epochs in all quarters of the world. Some countries, such as the islands of the South Pacific, the interior of Africa, and parts of North and South America, have passed direct from the Stone to the Iron Age. In Europe the Iron Age may be said to cover the last years of the prehistoric and the early years of the historic periods. In Egypt, Chaldaea, Assyria, China, it reaches far back, to perhaps 4000 years before the Christian era. In Africa, where there has been no Bronze Age, the use of iron succeeded immediately the use of stone. In the Black Pyramid of Abusir (VIth Dynasty), at least 3000B.C., Gaston Maspero found some pieces of iron, and in the funeral text of Pepi I. (about 3400B.C.) the metal is mentioned. The use of iron in northern Europe would seem to have been fairly general long before the invasion of Caesar. But iron was not in common use in Denmark until the end of the 1st centuryA.D.In the north of Russia and Siberia its introduction was even as late asA.D.800, while Ireland enters upon her Iron Age about the beginning of the 1st century. In Gaul, on the other hand, the Iron Age dates back some 800 yearsB.C.; while in Etruria the metal was known some six centuries earlier. Homer represents Greece as beginning her Iron Age twelve hundred years before our era. The knowledge of iron spread from the south to the north of Europe. In approaching the East from the north of Siberia or from the south of Greece and the Troad, the history of iron in each country eastward is relatively later; while a review of European countries from the north towards the south shows the latter becoming acquainted with the metal earlier than the former It is suggested that these facts support the theory that it is from Africa that iron first came into use. The finding of worked iron in the Great Pyramids seems to corroborate this view. The metal, however, is singularly scarce in collections of Egyptian antiquities. The explanation of this would seem to lie in the fact that the relics are in most cases the paraphernalia of tombs, the funereal vessels and vases, and iron being considered an impure metal by the ancient Egyptians it was never used in their manufacture of these or for any religious purposes. This idea of impurity would seem a further proof of the African origin of iron. It was attributed to Seth, the spirit of evil who according to Egyptian tradition governed the central deserts of Africa. The Iron Age in Europe is characterized by an elaboration of designs in weapons, implements and utensils. These are no longer cast but hammered into shape, and decoration is elaborate curvilinear rather than simplerectilinear, the forms and character of the ornamentation of the northern European weapons resembling in some respects Roman arms, while in others they are peculiar and evidently representative of northern art. The dead were buried in an extended position, while in the preceding Bronze Age cremation had been the rule.
See Lord Avebury,Prehistoric Times(1865; 1900); Sir J. Evans,Ancient Stone Implements(1897);Horae Ferales, or Studies in the Archaeology of Northern Nations, by Kemble (1863); Gaston C. C. Maspero,Guide du Musée de Boulaq, 296;Scotland in Pagan Times—The Iron Age, by Joseph Anderson (1883).
See Lord Avebury,Prehistoric Times(1865; 1900); Sir J. Evans,Ancient Stone Implements(1897);Horae Ferales, or Studies in the Archaeology of Northern Nations, by Kemble (1863); Gaston C. C. Maspero,Guide du Musée de Boulaq, 296;Scotland in Pagan Times—The Iron Age, by Joseph Anderson (1883).
IRON AND STEEL.11. Iron, the most abundant and the cheapest of the heavy metals, the strongest and most magnetic of known substances, is perhaps also the most indispensable of all save the air we breathe and the water we drink. For one kind of meat we could substitute another; wool could be replaced by cotton, silk or fur; were our common silicate glass gone, we could probably perfect and cheapen some other of the transparent solids; but even if the earth could be made to yield any substitute for the forty or fifty million tons of iron which we use each year for rails, wire, machinery, and structural purposes of many kinds, we could not replace either the steel of our cutting tools or the iron of our magnets, the basis of all commercial electricity. This usefulness iron owes in part, indeed, to its abundance, through which it has led us in the last few thousands of years to adapt our ways to itsproperties; but still in chief part first to the single qualities in which it excels, such as its strength, its magnetism, and the property which it alone has of being made at will extremely hard by sudden cooling and soft and extremely pliable by slow cooling; second, to the special combinations of useful properties in which it excels, such as its strength with its ready welding and shaping both hot and cold; and third, to the great variety of its properties. It is a very Proteus. It is extremely hard in our files and razors, and extremely soft in our horse-shoe nails, which in some countries the smith rejects unless he can bend them on his forehead; with iron we cut and shape iron. It is extremely magnetic and almost non-magnetic; as brittle as glass and almost as pliable and ductile as copper; extremely springy, and springless and dead; wonderfully strong, and very weak; conducting heat and electricity easily, and again offering great resistance to their passage; here welding readily, there incapable of welding; here very infusible, there melting with relative ease. The coincidence that so indispensable a thing should also be so abundant, that an iron-needing man should be set on an iron-cored globe, certainly suggests design. The indispensableness of such abundant things as air, water and light is readily explained by saying that their very abundance has evolved a creature dependent on them. But the indispensable qualities of iron did not shape man’s evolution, because its great usefulness did not arise until historic times, or even, as in case of magnetism, until modern times.
These variations in the properties of iron are brought about in part by corresponding variations in mechanical and thermal treatment, by which it is influenced profoundly, and in part by variations in the proportions of certain foreign elements which it contains; for, unlike most of the other metals, it is never used in the pure state. Indeed pure iron is a rare curiosity. Foremost among these elements is carbon, which iron inevitably absorbs from the fuel used in extracting it from its ores. So strong is the effect of carbon that the use to which the metal is put, and indeed its division into its two great classes, the malleable one, comprising steel and wrought iron, with less than 2.20% of carbon, and the unmalleable one, cast iron, with more than this quantity, are based on carbon-content. (See Table I.)
Table I.—General Classification of Iron and Steel according (1) to Carbon-Content and (2) to Presence or Absence of Inclosed Slag.
They may be either Bessemer, open-hearth, or crucible steel. Malleable cast iron also often belongs here.
Normal cast iron, ”washed” metal, and most “malleable cast iron” belong here.
2.Nomenclature.—Until about 1860 there were only three important classes of iron—wrought iron, steel and cast iron. The essential characteristic of wrought iron was its nearly complete freedom from carbon; that of steel was its moderate carbon-content (say between 0.30 and 2.2%), which, though great enough to confer the property of being rendered intensely hard and brittle by sudden cooling, yet was not so great but that the metal was malleable when cooled slowly; while that of cast iron was that it contained so much carbon as to be very brittle whether cooled quickly or slowly. This classification was based on carbon-content, or on the properties which it gave. Beyond this, wrought iron, and certain classes of steel which then were important, necessarily contained much slag or “cinder,” because they were made by welding together pasty particles of metal in a bath of slag, without subsequent fusion. But the best class of steel, crucible steel, was freed from slag by fusion in crucibles; hence its name, “cast steel.” Between 1860 and 1870 the invention of the Bessemer and open-hearth processes introduced a new class of iron to-day called “mild” or “low-carbon steel,” which lacked the essential property of steel, the hardening power, yet differed from the existing forms of wrought iron in freedom from slag, and from cast iron in being very malleable. Logically it was wrought iron, the essence of which was that it was (1) “iron” as distinguished from steel, and(2) malleable,i.e.capable of being “wrought.” This name did not please those interested in the new product, because existing wrought iron was a low-priced material. Instead of inventing a wholly new name for the wholly new product, they appropriated the name “steel,” because this was associated in the public mind with superiority. This they did with the excuse that the new product resembled one class of steel—cast steel—in being free from slag; and, after a period of protest, all acquiesced in calling it “steel,” which is now its firmly established name. The old varieties of wrought iron, steel and cast iron preserve their old names; the new class is called steel by main force. As a result, certain varieties, such as blister steel, are called “steel” solely because they have the hardening power, and others, such as low-carbon steel, solely because they are free from slag. But the former lack the essential quality, slaglessness, which makes the latter steel, and the latter lack the essential quality, the hardening power, which makes the former steel. “Steel” has come gradually to stand rather for excellence than for any specific quality. These anomalies, however confusing to the general reader, in fact cause no appreciable trouble to important makers or users of iron and steel, beyond forming an occasional side-issue in litigation.
3.Definitions.—Wrought ironis slag-bearing malleable iron, containing so little carbon (0.30% or less), or its equivalent, that it does not harden greatly when cooled suddenly.
Steelis iron which is malleable at least in some one range of temperature, and also is either (a) cast into an initially malleable mass, or (b) is capable of hardening greatly by sudden cooling, or (c) is both so cast and so capable of hardening. (Tungsten steel and certain classes of manganese steel are malleable only when red-hot.) Normal or carbon steel contains between 0.30 and 2.20% of carbon, enough to make it harden greatly when cooled suddenly, but not enough to prevent it from being usefully malleable when hot.
Cast ironis, generically, iron containing so much carbon (2.20% or more) or its equivalent that it is not usefully malleable at any temperature. Specifically, it is cast iron in the form of castings other than pigs, or remelted cast iron suitable for such castings, as distinguished from pig iron,i.e.the molten cast iron as it issues from the blast furnace, or the pigs into which it is cast.
Malleable cast ironis iron which has been cast in the condition of cast iron, and made malleable by subsequent treatment without fusion.
Alloy steelsandcast ironsare those which owe their properties chiefly to the presence of one or more elements other than carbon.
Ingot ironis slagless steel with less than 0.30% of carbon.
Ingot steelis slagless steel containing more than 0.30% of carbon.
Weld steelis slag-bearing iron malleable at least at some one temperature, and containing more than 0.30% of carbon.
4.Historical Sketch.—The iron oxide of which the ores of iron consist would be so easily deoxidized and thus brought to the metallic state by the carbon,i.e.by the glowing coals of any primeval savage’s wood fire, and the resulting metallic iron would then differ so strikingly from any object which he had previously seen, that its very early use by our race is only natural. The first observing savage who noticed it among his ashes might easily infer that it resulted from the action of burning wood on certain extremely heavy stones. He could pound it out into many useful shapes. The natural steps first of making it intentionally by putting such stones into his fire, and next of improving his fire by putting it and these stones into a cavity on the weather side of some bank with an opening towards the prevalent wind, would give a simple forge, differing only in size, in lacking forced blast, and in details of construction, from the Catalan forges and bloomaries of to-day. Moreover, the coals which deoxidized the iron would inevitably carburize some lumps of it, here so far as to turn it into the brittle and relatively useless cast iron, there only far enough to convert it into steel, strong and very useful even in its unhardened state. Thus it is almost certain that much of the earliest iron was in fact steel. How soon after man’s discovery, that he could beat iron and steel out while cold into useful shapes, he learned to forge it while hot is hard to conjecture. The pretty elaborate appliances, tongs or their equivalent, which would be needed to enable him to hold it conveniently while hot, could hardly have been devised till a very much later period; but then he may have been content to forge it inconveniently, because the great ease with which it mashes out when hot, perhaps pushed with a stout stick from the fire to a neighbouring flat stone, would compensate for much inconvenience. However this may be, very soon after man began to practise hot-forging he would inevitably learn that sudden cooling, by quenching in water, made a large proportion of his metal, his steel, extremely hard and brittle, because he would certainly try by this very quenching to avoid the inconvenience of having the hot metal about. But the invaluable and rather delicate art of tempering the hardened steel by a very careful and gentle reheating, which removes its extreme brittleness though leaving most of its precious hardness, needs such skilful handling that it can hardly have become known until very long after the art of hot-forging.
The oxide ores of copper would be deoxidized by the savage’s wood fire even more easily than those of iron, and the resulting copper would be recognized more easily than iron, because it would be likely to melt and run together into a mass conspicuous by its bright colour and its very great malleableness. From this we may infer that copper and iron probably came into use at about the same stage in man’s development, copper before iron in regions which had oxidized copper ores, whether they also had iron ores or not, iron before copper in places where there were pure and easily reduced ores of iron but none of copper. Moreover, the use of each metal must have originated in many different places independently. Even to-day isolated peoples are found with their own primitive iron-making, but ignorant of the use of copper.
If iron thus preceded copper in many places, still more must it have preceded bronze, an alloy of copper and tin much less likely than either iron or copper to be made unintentionally. Indeed, though iron ores abound in many places which have neither copper nor tin, yet there are but few places which have both copper and tin. It is not improbable that, once bronze became known, it might replace iron in a measure, perhaps even in a very large measure, because it is so fusible that it can be cast directly and easily into many useful shapes. It seems to be much more prominent than iron in the Homeric poems; but they tell us only of one region at one age. Even if a nation here or there should give up the use of iron completely, that all should is neither probable nor shown by the evidence. The absence of iron and the abundance of bronze in the relics of a prehistoric people is a piece of evidence to be accepted with caution, because the great defect of iron, its proneness to rust, would often lead to its complete disappearance, or conversion into an unrecognizable mass, even though tools of bronze originally laid down beside it might remain but little corroded. That the ancients should have discovered an art of hardening bronze is grossly improbable, first because it is not to be hardened by any simple process like the hardening of steel, and second because, if they had, then a large proportion of the ancient bronze tools now known ought to be hard, which is not the case.
Because iron would be so easily made by prehistoric and even by primeval man, and would be so useful to him, we are hardly surprised to read in Genesis that Tubal Cain, the sixth in descent from Adam, discovered it; that the Assyrians had knives and saws which, to be effective, must have been of hardened steel,i.e.of iron which had absorbed some carbon from the coals with which it had been made, and had been quenched in water from a red heat; that an iron tool has been found embedded in the ancient pyramid of Kephron (probably as early as 3500B.C.); that iron metallurgy had advanced at the time of Tethmosis (Thothmes) III. (about 1500B.C.) so far that bellows were used for forcing the forge fire; that in Homer’s time (not later than the 9th centuryB.C.) the delicate art of hardening and tempering steel was so familiar that the poet used it for a simile, likeningthe hissing of the stake which Ulysses drove into the eye of Polyphemus to that of the steel which the smith quenches in water, and closing with a reference to the strengthening effect of this quenching; and that at the time of Pliny (A.D.23-79) the relative value of different baths for hardening was known, and oil preferred for hardening small tools. These instances of the very early use of this metal, intrinsically at once so useful and so likely to disappear by rusting away, tell a story like that of the single foot-print of the savage which the waves left for Robinson Crusoe’s warning. Homer’s familiarity with the art of tempering could come only after centuries of the wide use of iron.
3.Three Periods.—The history of iron may for convenience be divided into three periods: a first in which only the direct extraction of wrought iron from the ore was practised; a second which added to this primitive art the extraction of iron in the form of carburized or cast iron, to be used either as such or for conversion into wrought iron; and a third in which the iron worker used a temperature high enough to melt wrought iron, which he then called molten steel. For brevity we may call these the periods of wrought iron, of cast iron, and of molten steel, recognizing that in the second and third the earlier processes continued in use. The first period began in extremely remote prehistoric times; the second in the 14th century; and the third with the invention of the Bessemer process in 1856.