1911 Encyclopædia Britannica/Iron

IRON [symbol Fe, atomic weight 55·85 (O＝16)], a metallic chemical element. Although iron occurs only sparingly in the free state, the abundance of ores from which it may be readily obtained led to its application in the arts at a very remote period. It is generally agreed, however, that the Iron Age, the period of civilization during which this metal played an all-important part, succeeded the ages of copper and bronze, notwithstanding the fact that the extraction of these metals required greater metallurgical skill. The Assyrians and Egyptians made considerable use of the metal; and in Genesis iv. 22 mention is made of Tubal-cain as the instructor of workers in iron and copper. The earlier sources of the ores appear to have been in India; the Greeks, however, obtained it from the Chalybes, who dwelt on the south coast of the Black Sea; and the Romans, besides drawing from these deposits, also exploited Spain, Elba and the province of Noricum. (See Metal-work.)

The chief occurrences of metallic iron are as minute spiculae disseminated through basaltic rocks, as at Giant’s Causeway and in the Auvergne, and, more particularly, in meteorites (q.v.). In combination it occurs, usually in small quantity, in most natural waters, in plants, and as a necessary constituent of blood. The economic sources are treated under Iron and Steel below; in the same place will be found accounts of the manufacture, properties, and uses of the metal, the present article being confined to its chemistry. The principal iron ores are the oxides and carbonates, and these readily yield the metal by smelting with carbon. The metal so obtained invariably contains a certain amount of carbon, free or combined, and the proportion and condition regulate the properties of the metal, giving origin to the three important varieties: cast iron, steel, wrought iron. The perfectly pure metal may be prepared by heating the oxide or oxalate in a current of hydrogen; when obtained at a low temperature it is a black powder which oxidizes in air with incandescence; produced at higher temperatures the metal is not pyrophoric. Péligot obtained it as minute tetragonal octahedra and cubes by reducing ferrous chloride in hydrogen. It may be obtained electrolytically from solutions of ferrous and magnesium sulphates and sodium bicarbonate, a wrought iron anode and a rotating cathode of copper, thinly silvered and iodized, being employed (S. Maximowitsch, Zeit. Elektrochem., 1905, 11, p. 52).

In bulk, the metal has a silvery white lustre and takes a high polish. Its specific gravity is 7.84; and the average specific heat over the range 15°–100° is 0.10983; this value increases with temperature to 850°, and then begins to diminish. It is the most tenacious of all the ductile metals at ordinary temperatures with the exception of cobalt and nickel; it becomes brittle, however, at the temperature of liquid air. It softens at a red heat, and may be readily welded at a white heat; above this point it becomes brittle. It fuses at about 1550°–1600°, and may be distilled in the electric furnace (H. Moissan, Compt. rend., 1906, 142, p. 425). It is attracted by a magnet and may be magnetized, but the magnetization is quickly lost. The variation of physical properties which attends iron on heating has led to the view that the metal exists in allotropic forms (see Iron and Steel, below).

Iron is very reactive chemically. Exposed to atmospheric influences it is more or less rapidly corroded, giving the familiar rust (q.v.). S. Burnie (Abst. J.C.S., 1907, ii. p. 469) has shown that water is decomposed at all temperatures from 0° to 100° by the finely divided metal with liberation of hydrogen, the action being accelerated when oxides are present. The decomposition of steam by passing it through a red-hot gun-barrel, resulting in the liberation of hydrogen and the production of magnetic iron oxide, Fe₃O₄, is a familiar laboratory method for preparing hydrogen (q.v.). When strongly heated iron inflames in oxygen and in sulphur vapour; it also combines directly with the halogens. It dissolves in most dilute acids with liberation of hydrogen; the reaction between sulphuric acid and iron turnings being used for the commercial manufacture of this gas. It dissolves in dilute cold nitric acid with the formation of ferrous and ammonium nitrates, no gases being liberated; when heated or with stronger acid ferric nitrate is formed with evolution of nitrogen oxides.

It was observed by James Keir (Phil. Trans., 1790, p. 359) that iron, after having been immersed in strong nitric acid, is insoluble in acids, neither does it precipitate metals from solutions. This “passivity” may be brought about by immersion in other solutions, especially by those containing such oxidizing anions as NO′₃, ClO′₃, less strongly by the anions SO″₄ CN′, CNS′, C₂H₃O′₂, OH′, while Cl′, Br′ practically inhibit passivity; H′ is the only cation which has any effect, and this tends to exclude passivity. It is also occasioned by anodic polarization of iron in sulphuric acid. Other metals may be rendered passive; for example, zinc does not precipitate copper from solutions of the double cyanides and sulphocyanides, nickel and cadmium from the nitrates, and iron from the sulphate, but it immediately throws down nickel and cadmium from the sulphates and chlorides, and lead and copper from the nitrates (see O. Sackur, Zeit. Elektrochem., 1904, 10, p. 841). Anodic polarization in potassium chloride solution renders molybdenum, niobium, ruthenium, tungsten, and vanadium passive (W. Muthmann and F. Frauenberger, Sitz. Bayer. Akad. Wiss., 1904, 34, p. 201), and also gold in commercial potassium cyanide solution (A. Coehn and C. L. Jacobsen, Abs. J.C.S., 1907, ii. p. 926). Several hypotheses have been promoted to explain this behaviour, and, although the question is not definitely settled, the more probable view is that it is caused by the formation of a film of an oxide, a suggestion made many years ago by Faraday (see P. Krassa, Zeit. Elektrochem., 1909, 15, p. 490). Fredenhagen (Zeit. physik. Chem., 1903, 43, p. 1), on the other hand, regarded it as due to surface films of a gas; submitting that the difference between iron made passive by nitric acid and by anodic polarization was explained by the film being of nitrogen oxides in the first case and of oxygen in the second case. H. L. Heathcote and others regard the passivity as invariably due to electrolytic action (see papers in the Zeit. physik. Chem., 1901 et seq.).

Compounds of Iron.

Oxides and Hydroxides.—Iron forms three oxides: ferrous oxide, FeO, ferric oxide, Fe₂O₃, and ferroso-ferric oxide, Fe₃O₄. The first two give origin to well-defined series of salts, the ferrous salts, wherein the metal is divalent, and the ferric salts, wherein the metal is trivalent; the former readily pass into the latter on oxidation, and the latter into the former on reduction.

Ferrous oxide is obtained when ferric oxide is reduced in hydrogen at 300° as a black pyrophoric powder. Sabatier and Senderens (Compt. rend., 1892, 114, p. 1429) obtained it by acting with nitrous oxide on metallic iron at 200°, and Tissandier by heating the metal to 900° in carbon dioxide; Donau (Monats., 1904, 25, p. 181), on the other hand, obtained a magnetic and crystalline-ferroso-ferric oxide at 1200°. It may also be prepared as a black velvety powder which readily takes up oxygen from the air by adding ferrous oxalate to boiling caustic potash. Ferrous hydrate, Fe(OH)₂, when prepared from a pure ferrous salt and caustic soda or potash free from air, is a white powder which may be preserved in an atmosphere of hydrogen. Usually, however, it forms a greenish mass, owing to partial oxidation. It oxidizes on exposure with considerable evolution of heat; it rapidly absorbs carbon dioxide; and readily dissolves in acids to form ferrous salts, which are usually white when anhydrous, but greenish when hydrated.

Ferric oxide or iron sesquioxide, Fe₂O₃, constitutes the valuable ores red haematite and specular iron; the minerals brown haematite or limonite, and göthite and also iron rust are hydrated forms. It is obtained as a steel-grey crystalline powder by igniting the oxide or any ferric salt containing a volatile acid. Small crystals are formed by passing ferric chloride vapour over heated lime. When finely ground these crystals yield a brownish red powder which dissolves slowly in acids, the most effective solvent being a boiling mixture of 8 parts of sulphuric acid and 3 of water. Ferric oxide is employed as a pigment, as jeweller’s rouge, and for polishing metals. It forms several hydrates, the medicinal value of which was recognized in very remote times. Two series of synthetic hydrates were recognized by Muck and Tommasi: the “red” hydrates, obtained by precipitating ferric salts with alkalis, and the “yellow” hydrates, obtained by oxidizing moist ferrous hydroxide or carbonates. J. van Bemmelen has shown that the red hydrates are really colloids, the amount of water retained being such that its vapour pressure equals the pressure of the aqueous vapour in the superincumbent atmosphere. By heating freshly prepared red ferric hydrate with water under 5000 atmospheres pressure Ruff (Ber., 1901, 34, p. 3417) obtained definite hydrates corresponding to the minerals limonite (30°-42.5°), göthite (42.5°–62.5°), and hydrohaematite (above 62.5°). Thomas Graham obtained a soluble hydrate by dissolving the freshly prepared hydrate in ferric chloride and dialysing the solution, the soluble hydrate being left in the dialyser. All the chlorine, however, does not appear to be removed by this process, the residue having the composition 82Fe(OH)₃·FeCl₃; but it may be by electrolysing in a porous cell (Tribot and Chrétien, Compt. rend., 1905, 140, p. 144). On standing, the solution usually gelatinizes, a process accelerated by the addition of an electrolyte. It is employed in medicine under the name Liquor ferri dialysati. The so-called soluble meta-ferric hydroxide, FeO(OH)(?), discovered by Péan de St Gilles in 1856, may be obtained by several methods. By heating solutions of certain iron salts for some time and then adding a little sulphuric acid it is precipitated as a brown powder. Black scales, which dissolve in water to form a red solution, are obtained by adding a trace of hydrochloric acid to a solution of basic ferric nitrate which has been heated to 100° for three days. A similar compound, which, however, dissolves in water to form an orange solution, results by adding salt to a heated solution of ferric chloride. These compounds are insoluble in concentrated, but dissolve readily in dilute acids.

Red ferric hydroxide dissolves in acids to form a well-defined series of salts, the ferric salts, also obtained by oxidizing ferrous salts; they are usually colourless when anhydrous, but yellow or brown when hydrated. It has also feebly acidic properties, forming ferrites with strong bases.

Magnetite, Fe₃O₄, may be regarded as ferrous ferrite, FeO·Fe₂O₃. This important ore of iron is most celebrated for its magnetic properties (see Magnetism and Compass), but the mineral is not always magnetic, although invariably attracted by a magnet. It may be obtained artificially by passing steam over red-hot iron. It dissolves in acids to form a mixture of a ferrous and ferric salt,^[1] and if an alkali is added to the solution a black precipitate is obtained which dries to a dark brown mass of the composition Fe(OH)₂·Fe₂O₃; this substance is attracted by a magnet, and thus may be separated from the admixed ferric oxide. Calcium ferrite, magnesium ferrite and zinc ferrite, RO·Fe₂O₃ (R = Ca, Mg, Zn), are obtained by intensely heating mixtures of the oxides; magnesium ferrite occurs in nature as the mineral magnoferrite, and zinc ferrite as franklinite, both forming black octahedra.

Ferric acid, H₂FeO₄. By fusing iron with saltpetre and extracting the melt with water, or by adding a solution of ferric nitrate in nitric acid to strong potash, an amethyst or purple-red solution is obtained which contains potassium ferrate. E. Frémy investigated this discovery, made by Stahl in 1702, and showed that the same solution resulted when chlorine is passed into strong potash solution containing ferric hydrate in suspension. Haber and Pick (Zeit. Elektrochem., 1900, 7, p. 215) have prepared potassium ferrate by electrolysing concentrated potash solution, using an iron anode. A temperature of 70°, and a reversal of the current (of low density) between two cast iron electrodes every few minutes, are the best working conditions. When concentrated the solution is nearly black, and on heating it yields a yellow solution of potassium ferrite, oxygen being evolved. Barium ferrate, BaFeO₄·H₂O, obtained as a dark red powder by adding barium chloride to a solution of potassium ferrate, is fairly stable. It dissolves in acetic acid to form a red solution, is not decomposed by cold sulphuric acid, but with hydrochloric or nitric acid it yields barium and ferric salts, with evolution of chlorine or oxygen (Baschieri, Gazetta, 1906, 36, ii. p. 282).

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 41/2H₂O) 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 as oleum 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 (see Solution). 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 name Liquor 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 Fe₇S₈ and Fe₈S₉, i.e. 5FeS·Fe₂S₃ and 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 the Tractatus chymico-philosophicus ascribed 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 amounts of 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. Fe₂P₃ is 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)₃ and 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 under Iron 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 see Prussic 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 also Chemistry, § Analytical). For the quantitative estimation see Assaying.

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 preparation Vinum 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 is Trochiscus ferri redacti (strength, 1 grain of reduced iron in each).

3. Ferri sulphas, ferrous sulphate, from which is prepared Mistura 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, 4/5 gr. of quinine sulphate, and 1/32 gr. 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 perchloridi and Tinctura 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 prepared Vinum 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 21/2% 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 of post partum haemorrhage. 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.

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1. By 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.