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
