[7 bis] The most important of the compounds corresponding with chromic oxide is chromic chloride, Cr2Cl6, which is known in an anhydrous and in a hydrated form. It resembles ferric and aluminic chlorides in many respects. There is a great difference between the anhydrous and the hydrated chlorides; the former is insoluble in water, the latter easily dissolves, and on evaporation its solution forms a hygroscopic mass which is very unstable and easily evolves hydrochloric acid when heated with water. The anhydrous form is of a violet colour, and Wöhler gives the following method for its preparation: an intimate mixture is prepared of the anhydrous chromic oxide with carbon and organic matter, and charged into a wide infusible glass or porcelain tube which is heated in a combustion furnace; one extremity of the tube communicates with an apparatus generating chlorine which is passed through several bottles containing sulphuric acid in order to dry it perfectly before it reaches the tube. On heating the portion of the tube in which the mixture is placed and passing chlorine through, a slightly volatile sublimate of chromic chloride, CrCl3 or Cr2Cl6, is formed. This substance forms violet tabular crystals, which may be distilled in dry chlorine without change, but which, however, require a red heat for their volatilisation. These crystals are greasy to the touch and insoluble in water, but if they be powdered and boiled in water for a long time they pass into a green solution. Strong sulphuric acid does not act on the anhydrous salt, or only acts with exceeding slowness, like water. Even aqua regia and other acids do not act on the crystals, and alkalis only show a very feeble action. The specific gravity of the crystals is 2·99. When fused with sodium carbonate and nitre they give sodium chloride and potassium chromate, and when ignited in air they form green chromic oxide and evolve chlorine. On ignition in a stream of ammonia, chromic chloride forms sal-ammoniac and chromium nitride, CrN (analogous to the nitrides BN, AlN). Mosberg and Peligot showed that when chromic chloride is ignited in hydrogen, it parts with one-third of its chlorine, forming chromous chloride, CrCl2—that is, there is formed from a compound corresponding with chromic oxide, Cr2O3, a compound answering to the suboxide, chromous oxide, CrO—just as hydrogen converts ferric chloride into ferrous chloride with the aid of heat. Chromous chloride, CrCl2, forms colourless crystals easily soluble in water, which in dissolving evolve a considerable amount of heat, and form a blue liquid, capable of absorbing oxygen from the air with great facility, being converted thereby into a chromic compound.
The blue solution of chromous chloride may also be obtained by the action of metallic zinc on the green solution of the hydrated chromic chloride; the zinc in this case takes up chlorine just as the hydrogen did. It must be employed in a large excess. Chromic oxide is also formed in the action of zinc on chromic chloride, and if the solution remain for a long time in contact with the zinc the whole of the chromium is converted into chromic oxychloride. Other chromic salts are also reduced by zinc into chromous salts, CrX2, just as the ferric salts FeX3 are converted into ferrous salts FeX2 by it. The chromous salts are exceedingly unstable and easily oxidise and pass into chromic salts; hence the reducing power of these salts is very great. From cupric salts they separate cuprous salts, from stannous salts they precipitate metallic tin, they reduce mercuric salts into mercurous and ferric into ferrous salts. Moreover, they absorb oxygen from the air directly. With potassium chromate they give a brown precipitate of chromium dioxide or of chromic oxide, according to the relative amounts of the substances taken: CrO3 + CrO = 2CrO2 or CrO3 + 3CrO = 2Cr2O3. Aqueous ammonia gives a blue precipitate, and in the presence of ammoniacal salts a blue liquid is obtained which turns red in the air from oxidation. This is accompanied by the formation of compounds analogous to those given by cobalt (Chapter [XXII.]). A solution of chromous chloride with a hot saturated solution of sodium acetate, C2H3NaO2, gives, on cooling, transparent red crystals of chromous acetate, C4H6CrO4,H2O. This salt is also a powerful reducing agent, but may be kept for a long time in a vessel full of carbonic anhydride.
The insoluble anhydrous chromic chloride CrCl3 very easily passes into solution in the presence of a trace (0·004) of chromous chloride CrCl2. This remarkable phenomenon was observed by Peligot and explained by Löwel in the following manner: chromous chloride, as a lower stage of oxidation, is capable of absorbing both oxygen and chlorine, combining with various substances. It is able to decompose many chlorides by taking up chlorine from them; thus it precipitates mercurous chloride from a solution of mercuric chloride, and in so doing passes into chromic chloride: 2CrCl2 + 2HgCl2 = Cr2Cl6 + 2HgCl. Let us suppose that the same phenomenon takes place when the anhydrous chromic chloride is mixed with a solution of chromous chloride. The latter will then take up a portion of the chlorine of the former, and pass into a soluble hydrate of chromic chloride (hydrochloride of oxide of chromium), and the original anhydrous chromic chloride will pass into chromous chloride. The chromous chloride re-formed in this manner will then act on a fresh quantity of the chromic chloride, and in this manner transfer it entirely into solution as hydrate. This view is confirmed by the fact that other chlorides, capable of absorbing chlorine like chromous chloride, also induce the solution of the insoluble chromic chloride—for example, ferrous chloride, FeCl2, and cuprous chloride. The presence of zinc also aids the solution of chromic chloride, owing to its converting a portion of it into chromous chloride. The solution of chromic chloride in water obtained by these methods is perfectly identical with that which is formed by dissolving chromic hydroxide in hydrochloric acid. On evaporating the green solution obtained in this manner, it gives a green mass, containing water. On further heating it leaves a soluble chromic oxychloride, and when ignited it first forms an insoluble oxychloride and then chromic oxide; but no anhydrous chromic chloride, Cr2Cl6, is formed by heating the aqueous solution of chromic chloride, which forms an important fact in support of the view that the green solution of chromic chloride is nothing else but hydrochloride of oxide of chromium. At 100° the composition of the green hydrate is Cr2Cl6,9H2O, and on evaporation at the ordinary temperature over H2SO4 crystals are obtained with 12 equivalents of water; the red mass obtained at 120° contains Cr2O3,4Cr2Cl6,24H2O. The greater portion of it is soluble in water, like the mass which is formed at 150°. The latter contains Cr2O3,2Cr2Cl6,9H2O = 3(Cr2OCl4,3H2O)—that is, it presents the same composition as chromic chloride in which one atom of oxygen replaces two of chlorine. And if the hydrate of chromic chloride be regarded as Cr2O3,6HCl, the substance which is obtained should be regarded as Cr2O3,4HCl combined with water, H2O. The addition of alkalis—for example, baryta—to a solution of chromic chloride immediately produces a precipitate, which, however, re-dissolves on shaking, owing to the formation of one of the oxychlorides just mentioned, which may be regarded as basic salts. Thus we may represent the product of the change produced on chromic chloride under the influence of water and heat by the following formulæ: first Cr2O3,6HCl or Cr2Cl6,3H2O is formed, then Cr2O3,4HCl,H2O or Cr2OCl4,3H2O, and lastly Cr2O3,2HCl,2H2O or Cr2O2Cl2,3H2O. In all three cases there are 2 equivalents of chromium to at least 3 equivalents of water. These compounds may be regarded as being intermediate between chromic hydroxide and chloride; chromic chloride is Cr2Cl6, the first oxychloride Cr2(OH)2Cl4, the second Cr2(OH)4Cl2, and the hydrate Cr2(OH)6—that is, the chlorine is replaced by hydroxyl.
It is very important to remark two circumstances in respect to this: (1) That the whole of the chlorine in the above compounds is not precipitated from their solutions by silver nitrate; thus the normal salt of the composition Cr2Cl6,9H2O only gives up two-thirds of its chlorine; therefore Peligot supposes that the normal salt contains the oxychloride combined with hydrochloric acid: Cr2Cl6 + 2H2O = Cr2O2Cl2,4HCl, and that the chlorine held as hydrochloric acid reacts with the silver, whilst that held in the oxychloride does not enter into reaction, just as we observe a very feebly-developed faculty for reaction in the anhydrous chromic chloride; and (2) if the green aqueous solution of CrCl3 be left to stand for some time, it ultimately turns violet; in this form the whole of the chlorine is precipitated by AgNO3, whilst boiling re-converts it into the green variety. Löwel obtained the violet solution of hydrochloride of chromic oxide by decomposing the violet chromic sulphate with barium chloride. Silver nitrate precipitates all the chlorine from this violet modification; but if the violet solution be boiled and so converted into the green modification, silver nitrate then only precipitates a portion of the chlorine.
Recoura (1890–1893) obtained a crystallohydrate of violet chromium sulphate, Cr2(SO4)3, with 18 or 15 H2O. By boiling a solution of this crystallohydrate, he converted it into the green salt, which, when treated with alkalis, gave a precipitate of Cr2O3,2H2O, soluble in 2H2SO4 (and not 3), and only forming the basic salt, Cr2(OH)2(SO4)2. He therefore concludes that the green salts are basic salts. The cryoscopic determinations made by A. Speransky (1892) and Marchetti (1892) give a greater ‘depression’ for the violet than the green salts, that is, indicate a greater molecular weight for the green salts. But as Étard, by heating the violet sulphate to 100°, converted it into a green salt of the same composition, but with a smaller amount of H2O, it follows that the formation of a basic salt alone is insufficient to explain the difference between the green and violet varieties, and this is also shown by the fact that BaCl2 precipitates the whole of the sulphuric acid of the violet salt, and only a portion of that of the green salt. A. Speransky also showed that the molecular electro-conductivity of the green solutions is less than that of the violet. It is also known that the passage of the former into the latter is accompanied by an increase of volume, and, according to Recoura, by an evolution of heat also.
Piccini's researches (1894) throw an important light upon the peculiarities of the green chromium trichloride (or chromic chloride); he showed (1) that AgF (in contradistinction to the other salts of silver) precipitates all the chlorine from an aqueous solution of the green variety; (2) that solutions of green CrCl3,6H2O in ethyl alcohol and acetone precipitate all their chlorine when mixed with a similar solution of AgNO3; (3) that the rise of the boiling-point of the ethyl alcohol and acetone green solutions of CrCl3,6H2O (Chapter VII., Note [27 bis]) shows that i in this case (as in the aqueous solutions of MgSO4 and HgCl2) is nearly equal to 1, that is, that they are like solutions of non-conductors; (4) that a solution of green CrCl3 in methyl alcohol at first precipitates about ⅞ of its chlorine (an aqueous solution about ⅔) when treated with AgNO3, but after a time the whole of the chlorine is precipitated; and (5) that an aqueous solution of the green variety gradually passes into the violet, while a methyl alcoholic solution preserves its green colour, both of itself and also after the whole of the chlorine has been precipitated by AgNO3. If, however, in an aqueous or methyl alcoholic solution only a portion of the chlorine be precipitated, the solution gradually turns violet. In my opinion the general meaning of all these observations requires further elucidation and explanation, which should be in harmony with the theory of solutions. Recoura, moreover, obtained compounds of the green salt, Cr2(SO4)3, with 1, 2, and 3 molecules of H2SO4, K2SO4, and even a compound Cr2(SO4)3H2CrO4. By neutralising the sulphuric acid of the compounds of Cr2(SO4)3 and H2SO4 with caustic soda, Recoura obtained an evolution of 33 thousand calories per each 2NaHO, while free H2SO4 only gives 30·8 thousand calories. Recoura is of opinion that special chromo sulphuric acids, for instance (CrSO4)H2SO4 = ½Cr2(SO4)3H2SO4, are formed. With a still larger excess of sulphuric acid, Recoura obtained salts containing a still greater number of sulphuric acid radicles, but even this method does not explain the difference between the green and violet salts.
These facts must naturally be taken into consideration in order to arrive at any complete decision as to the cause of the different modifications of the chromic salts. We may observe that the green modification of chromic chloride does not give double salts with the metallic chlorides, whilst the violet variety forms compounds Cr2Cl6,2RCl (where R = an alkali metal), which are obtained by heating the chromates with an excess of hydrochloric acid and evaporating the solution until it acquires a violet colour. As the result of all the existing researches on the green and violet chromic salts, it appears to me most probable that their difference is determined by the feeble basic character of chromic oxide, by its faculty of giving basic salts, and by the colloidal properties of its hydroxide (these three properties are mutually connected), and moreover, it seems to me that the relation between the green and violet salts of chromic oxide best answers to the relation of the purpureo to the luteo cobaltic salts (Chapter XXII., Note [35]). This subject cannot yet be considered as exhausted (see Note [7]).
We may here observe that with tin the chromic salts, CrX3, give at low temperatures CrX2 and SnX2, whilst at high temperatures, on the contrary, CrX2 reduces the metal from its salts SnX2. The reaction, therefore, belongs to the class of reversible reactions (Beketoff).
Poulenc obtained anhydrous CrF3 (sp. gr. 3·78) and CrF2 (sp. gr. 4·11) by the action of gaseous HF upon CrCl2. A solution of fluoride of chromium is employed as a mordant in dyeing. Recoura (1890) obtained green and violet varieties of Cr2Br6,6H2O. The green variety can only be kept in the presence of an excess of HBr in the solution; if alone its solution easily passes into the violet variety with evolution of heat.
[8] The reduction of metallic chromium proceeds with comparative ease in aqueous solutions. Thus the action of sodium amalgams upon a strong solution of Cr2Cl6 gives (first CrCl2) an amalgam of chromium from which the mercury may be easily driven off by heating (in hydrogen to avoid oxidation), and there remains a spongy mass of easily oxidizable chromium. Plaset (1891), by passing an electric current through a solution of chrome alum mixed with a small amount of H2SO4 and K2SO4, obtained hard scales of chromium of a bluish-white colour possessing great hardness and stability (under the action of water, air, and acids). Glatzel (1890) reduced a mixture of 2KCl + Cr2Cl6 by heating it to redness with shavings of magnesium. The metallic chromium thus obtained has the appearance of a fine light-grey powder which is seen to be crystalline under the microscope; its sp. gr. at 16° is 6·7284. It fuses (with anhydrous borax) only at the highest temperatures, and after fusion presents a silver-white fracture. The strongest magnet has no action upon it.