Iron readily combines with non-metals—for instance, with chlorine, iodine, bromine, sulphur, and even with phosphorus and carbon; but on the other hand the property of combining with metals is but little developed in it—that is to say, it does not easily form alloys. Mercury, which acts on most metals, does not act directly on iron, and the iron amalgam, or solution of iron in mercury, which is used for electrical machines, is only obtained in a particular way—namely, with the co-operation of a sodium amalgam, in which the iron dissolves and by means of which it is reduced from solutions of its salts.
When iron acts on acids it forms ferrous salts of the type FeX2, and in the presence of air and oxidising agents they change by degrees into ferric salts of the type FeX3. This faculty of passing from the ferrous to the ferric state is still further developed in ferrous hydroxide. If sodium hydroxide be added to a solution of ferrous sulphate or green vitriol, FeSO4,[17] a white precipitate of ferrous hydroxide, FeH2O2, is obtained; but on exposure to the air, even under water, it turns green, becomes grey, and finally turns brown, which is due to the oxidation that it undergoes. Ferrous hydroxide is very sparingly soluble in water; the solution has, however, a distinct alkaline reaction, which is due to its being a fairly energetic basic oxide. In any case, ferrous oxide is far more energetic than ferric oxide, so that if ammonia be added to a solution containing a mixture of a ferrous and ferric salt, at first ferric hydroxide only will be precipitated. If barium carbonate, BaCO3, be shaken up in the cold with ferrous salts, it does not precipitate them—that is, does not change them into ferrous carbonate; but it completely separates all the iron from the ferric salts in the cold, according to the equation Fe2Cl6 + 3BaCO3 + 3H2O = Fe2O3,3H2O + 3BaCl2 + 3CO2. If ferrous hydroxide be boiled with a solution of potash, the water is decomposed, hydrogen is evolved, and the ferrous hydroxide is oxidised. The ferrous salts are in all respects similar to the salts of magnesium and zinc; they are isomorphous with them, but differ from them in that the ferrous hydroxide is not soluble either in aqueous potash or ammonia. In the presence of an excess of ammonium salts, however, a certain proportion of the iron is not precipitated by alkalis and alkali carbonates, which fact points to the formation of double ammonium salts.[18] The ferrous salts have a dull greenish colour, and form solutions also of a pale green colour, whilst the ferric salts have a brown or reddish-brown colour. The ferrous salts, being capable of oxidation, form very active reducing agents—for instance, under their action gold chloride, AuCl3, deposits metallic gold, nitric acid is transformed into lower oxides, and the highest oxides of manganese also pass into the lower forms of oxidation. All these reactions take place with especial ease in the presence of an excess of acid. This depends on the fact that the ferrous oxide, FeO (or salt), acting as a reducing agent, turns into ferric oxide, Fe2O3 (or salt), and in the ferric state it requires more acid for the formation of a normal salt than in the ferrous condition. Thus in the normal ferrous sulphate, FeSO4, there is one equivalent of iron to one equivalent of sulphur (in the sulphuric radicle), but in the neutral ferric salt, Fe2(SO4)3, there is one equivalent of iron to one and a half of sulphur in the form of the elements of sulphuric acid.[19]
The most simple oxidising agent for transforming ferrous into ferric salts is chlorine in the presence of water—for instance, 2FeCl2 + Cl2 = Fe2Cl6, or, generally speaking, 2FeO + Cl2 + H2O = Fe2O3 + 2HCl. When such a transformation is required it is best to add potassium chlorate and hydrochloric acid to the ferrous solution; chlorine is formed by their mutual reaction and acts as an oxidising agent. Nitric acid produces a similar effect, although more slowly. Ferrous salts may be completely and rapidly oxidised into ferric salts by means of chromic acid or permanganic acid, HMnO4, in the presence of acids—for example, 10FeSO4 + 2KMnO4 + 8H2SO4 = 5Fe2(SO4)3 + 2MnSO4 + K2SO4 + 8H2O. This reaction is easily observed by the change of colour, and its termination is easily seen, because potassium permanganate forms solutions of a bright red colour, and when added to a solution of a ferrous salt the above reaction immediately takes place in the presence of acid, and the solution then becomes colourless, because all the substances formed are only faintly coloured in solution. Directly all the ferrous compound has passed into the ferric state, any excess of permanganate which is added communicates a red colour to the liquid (see Chapter [XXI.])
Thus when ferrous salts are acted on by oxidising agents, they pass into the ferric form, and under the action of reducing agents the reverse reaction occurs. Sulphuretted hydrogen may, for instance, be used for this complete transformation, for under its influence ferric salts are reduced with separation of sulphur—for example, Fe2Cl6 + H2S = 2FeCl2 + 2HCl + S. Sodium thiosulphate acts in a similar way: Fe2Cl6 + Na2S2O3 + H2O = 2FeCl2 + Na2SO4 + 2HCl + S. Metallic iron or zinc,[20] in the presence, of acids, or sodium amalgam, &c., acts like hydrogen, and has also a similar reducing action, and this furnishes the best method for reducing ferric salts to ferrous salts—for instance, Fe2Cl6 + Zn = 2FeCl2 + ZnCl2. Thus the transition from ferrous salts to ferric salts and vice versâ is always possible.[21]
Ferric oxide, or sesquioxide of iron, Fe2O3, is found in nature, and is artificially prepared in the form of a red powder by many methods. Thus after heating green vitriol a red oxide of iron remains, called colcothar, which is used as an oil paint, principally for painting wood. The same substance in the form of a very fine powder (rouge) is used for polishing glass, steel, and other objects. If a mixture of ferrous sulphate with an excess of common salt be strongly heated, crystalline ferric oxide will be formed, having a dark violet colour, and resembling some natural varieties of this substance. When iron pyrites is heated for preparing sulphurous anhydride, ferric oxide also remains behind; it is used as a pigment. On the addition of alkalis to a solution of ferric salts, a brown precipitate of ferric hydroxide is formed, which when heated (even when boiled in water, that is, at about 100°, according to Tomassi) easily parts with the water, and leaves red anhydrous ferric oxide. Pure ferric oxide does not show any magnetic properties, but when heated to a white heat it loses oxygen and is converted into the magnetic oxide. Anhydrous ferric oxide which has been heated to a high temperature is with difficulty soluble in acids (but it is soluble when heated in strong acids, and also when fused with potassium hydrogen sulphate), whilst ferric hydroxide, at all events that which is precipitated from salts by means of alkalis, is very readily soluble in acids. The precipitated ferric hydroxide has the composition 2Fe2O33H2O, or Fe4H6O9. If this ordinary hydroxide be rendered anhydrous (at 100°), at a certain moment it becomes incandescent—that is, loses a certain quantity of heat. This self-incandescence depends on internal displacement produced by the transition of the easily-soluble (in acids) variety into the difficultly-soluble variety, and does not depend on the loss of water, since the anhydrous oxide undergoes the same change. In addition to this there exists a ferric hydroxide, or hydrated oxide of iron, which, like the strongly-heated anhydrous iron oxide, is difficultly soluble in acids. This hydroxide on losing water, or after the loss of water, does not undergo such self-incandescence, because no such state of internal displacement occurs (loss of energy or heat) with it as that which is peculiar to the ordinary oxide of iron. The ferric hydroxide which is difficultly soluble in acids has the composition Fe2O3,H2O. This hydroxide is obtained by a prolonged ebullition of water in which ferric hydroxide prepared by the oxidation of ferrous oxide is suspended, and also sometimes by similar treatment of the ordinary hydroxide after it has been for a long time in contact with water. The transition of one hydroxide to another is apparent by a change of colour; the easily-soluble hydroxide is redder, and the sparingly-soluble hydroxide more yellow in colour.[22]
The normal salts of the composition Fe2X6 or FeX3 correspond with ferric oxide—for example, the exceedingly volatile ferric chloride, Fe2Cl6, which is easily prepared in the anhydrous state by the action of chlorine on heated iron.[23] Such also is the normal ferric nitrate, Fe2(NO3)6; it is obtained by dissolving iron in an excess of nitric acid, taking care as far as possible to prevent any rise of temperature.[24] The normal salt separates from the brown solution when it is concentrated under a bell jar over sulphuric acid. This salt, Fe2(NO3)6,9H2O, then crystallises in well-formed and perfectly colourless crystals,[25] which deliquesce in air, melt at 35°, and are soluble in and decomposed by water. The decomposition may be seen from the fact that the solution is brown and does not yield the whole of the salt again, but gives partly basic salt. The normal salt (only stable in the presence of an excess of HNO3) is completely decomposed with great facility by heating with water, even at 130°, and this is made use of for removing iron (and also certain other oxides of the form R2O3) from many other bases (of the form RO) whose nitrates are far more stable. The ferric salts, FeX3, in passing into ferrous salts, act as oxidising agents, as is seen from the fact that they not only liberate S from SH2, but also iodine from KI like many oxidising agents.[25 bis]
Iron forms one other oxide besides the ferric and ferrous oxides; this contains twice as much oxygen as the former, but is so very unstable that it can neither be obtained in the free state nor as a hydrate. Whenever such conditions of double decomposition occur as should allow of its separation in the free state, it decomposes into oxygen and ferric oxide. It is known in the state of salts, and is only stable in the presence of alkalis, and forms salts with them which have a decidedly alkaline reaction; it is therefore a feebly acid oxide. Thus when small pieces of iron are heated with nitre or potassium chlorate a potassium salt of the composition K2FeO4 is formed, and therefore the hydrate corresponding with this salt should have the composition H2FeO4. It is called ferric acid. Its anhydride ought to contain FeO3 or Fe2O6—twice as much oxygen as ferric oxide. If a solution of potassium ferrate be mixed with acid, the free hydrate ought to be formed, but it immediately decomposes (2K2FeO4 + 5H2SO4 = 2K2SO4 + Fe2(SO4)3 + 5H2O + O3), oxygen being evolved. If a small quantity of acid be taken, or if a solution of potassium ferrate be heated with solutions of other metallic salts, ferric oxide is separated—for instance:
2CuSO4 + 2K2FeO4 = 2K2SO4 + O3 + Fe2O3 + 2CuO.