resembles kneading dough, and the process introduced in England
       became known as puddling. It is evident that the puddled mass, or
       bloom, is a heterogeneous substance obtained by mixing, and hence
       one part of the mass will still be rich in carbon, another will be
       poor, some parts will contain oxide not reduced, &c. The further
       treatment of the puddled mass consists in hammering and drawing it
       out into flat pieces, which on being hammered become more
       homogeneous, and when several pieces are welded together and again
       hammered out a still more homogeneous mass is obtained. The
       quality of the steel and iron thus formed depends principally on
       their uniformity. The want of uniformity depends on the oxides
       remaining inside the mass, and on the variable distribution of the
       carbon throughout the mass. In order to obtain a more homogeneous
       metal for manufacturing articles out of steel, it is drawn into
       thin rods, which are tied together in bundles and then again
       hammered out. As an example of what may be attained in this
       direction, imitation Damascus steel may be cited; it consists of
       twisted and plaited wire, which is then hammered into a dense
       mass. (Real damascened wootz steel may be made by melting a
       mixture of the best iron with graphite (1/12) and iron rust; the
       article is then corroded with acid, and the carbon remains in the
       form of a pattern.)

       Steel and wrought iron are manufactured from cast iron by
       puddling. They are, however, obtained not only by this method but
       also by the _bloomery process_, which is carried out in a fire
       similar to a blacksmith's forge, fed with charcoal and provided
       with a blast; a pig of cast iron is gradually pushed into the
       fire, and portions of it melt and fall to the bottom of the
       hearth, coming into contact with an air blast, and are thus
       oxidised. The bloom thus formed is then squeezed and hammered. It
       is evident that this process is only available when the charcoal
       used in the fire does not contain any foreign matter which might
       injure the quality of the iron or steel--for instance, sulphur or
       phosphorus--and therefore only wood charcoal may be used with
       impunity, from which it follows that this process can only be
       carried on where the manufacture of iron can be conducted with
       this fuel. Coal and coke contain the above-mentioned impurities,
       and would therefore produce iron of a brittle nature, and thus it
       would be necessary to have recourse to puddling, where the fuel is
       burnt on a special hearth, separate from the cast iron, whereby
       the impurities of the fuel do not come into contact with it. The
       manufacture of steel from cast iron may also be conducted in
       fires; but, in addition to this, it is also now prepared by many
       other methods. One of the long-known processes is called
       _cementation_, by which steel is prepared from wrought iron but
       not from cast iron. For this process strips of iron are heated
       red-hot for a considerable time whilst immersed in powdered
       charcoal; during this operation the iron at the surface combines
       with the charcoal, which however does not penetrate; after this
       the iron strips are re-forged, drawn out again, and cemented anew,
       repeating this process until a steel of the desired quality is
       formed--that is, containing the requisite proportion of carbon.
       The _Bessemer_ process occupies the front rank among the newer
       methods (since 1856); it is so called from the name of its
       inventor. This process consists in running melted cast iron into
       converters (holding about 6 tons of cast iron)--that is,
       egg-shaped receivers, fig. 94, capable of revolving on trunnions
       (in order to charge in the cast iron and discharge the steel), and
       forcing a stream of air through small apertures at a considerable
       pressure. Combustion of the iron and carbon at an elevated
       temperature then takes place, resulting from the bubbles of oxygen
       thus penetrating the mass of the cast iron. The carbon, however,
       burns to a greater extent than the iron, and therefore a mass is
       obtained which is much poorer in carbon than cast iron. As the
       combustion proceeds very rapidly in the mass of metal, the
       temperature rises to such an extent that even the wrought iron
       which may be formed remains in a molten condition, whilst the
       steel, being more fusible than the wrought iron, remains very
       liquid. In half an hour the mass is ready. The purest possible
       cast iron is used in the Bessemer process, because sulphur and
       phosphorus do not burn out like carbon, silicon, and manganese.

       [Illustration: FIG. 94.--Bessemer converter, constructed of iron
       plate and lined with ganister. The air is carried by the tubes, L,
       O, D to the bottom, M, from which it passes by a number of holes
       into the converter. The converter is rotated on the trunnion _d_
       by means of the rack and pinion H, when it is required either to
       receive molten cast iron from the melting furnaces or to pour out
       the steel.]

       The presence of manganese enables the sulphur to be removed with
       the slag, and the presence of lime or magnesia, which are
       introduced into the lining of the converter, facilitates the
       removal of the phosphorus. This basic Bessemer process, or _Thomas
       Gilchrist process_, introduced about 1880, enables ores containing
       a considerable amount of phosphorus, which had hitherto only been
       used for cast iron, to be used for making wrought iron and steel.
       Naturally the greatest uniformity will be obtained by re-melting
       the metal. Steel is re-melted in small wind furnaces, in masses
       not exceeding 30 kilos; a liquid metal is formed, which may be
       cast in moulds. A mixture of wrought and cast iron is often used
       for making cast steel (the addition of a small amount of metallic
       Al improves the homogeneity of the castings, by facilitating the
       passage of the impurities into slag). Large steel castings are
       made by simultaneous fusion in several furnaces and crucibles; in
       this way, castings up to 80 tons or more, such as large ordnance,
       may be made. This molten, and therefore homogeneous, steel is
       called _cast steel_. Of late years the _Martin's process_ for the
       manufacture of steel has come largely into use; it was invented in
       France about 1860, and with the use of regenerative furnaces it
       enables large quantities of cast steel to be made at a time. It is
       based on the melting of cast iron with iron oxides and iron
       itself--for instance, pure ores, scrap, &c. There the carbon of
       the cast iron and the oxygen of the oxide form carbonic oxide, and
       the carbon therefore burns out, and thus cast steel is obtained
       from cast iron, providing, naturally, that there is a requisite
       proportion and corresponding degree of heat. The advantage of this
       process is that not only do the carbon, silicon, and manganese,
       but also a great part of the sulphur and phosphorus of the cast
       iron burn out at the expense of the oxygen of the iron oxides.
       During the last decade the manufacture of steel and its
       application for rails, armour plate, guns, boilers, &c., has
       developed to an enormous extent, thanks to the invention of cheap
       processes for the manufacture of large masses of homogeneous cast
       steel. Wrought iron may also be melted, but the heat of a blast
       furnace is insufficient for this. It easily melts in the
       oxyhydrogen flame. It may be obtained in a molten state directly
       from cast iron, if the latter be melted with nitre and
       sufficiently stirred up. Considerable oxidation then takes place
       inside the mass of cast iron, and the temperature rises to such an
       extent that the wrought iron formed remains liquid. A method is
       also known for obtaining wrought iron directly from rich iron ores
       by the action of carbonic oxide: the wrought iron is then formed
       as a spongy mass (which forms an excellent filter for purifying
       water), and may be worked up into wrought iron or steel either by
       forging or by dissolving in molten cast iron.

       Everybody is more or less familiar with the _difference in the
       properties of steel and wrought iron_. Iron is remarkable for its
       softness, pliability, and small elasticity, whilst steel may be
       characterised by its capability of attaining elasticity and
       hardness if it be cooled suddenly after having been heated to a
       definite temperature, or, as it is termed, _tempered_. But if
       tempered steel be re-heated and slowly cooled, it becomes as soft
       as wrought iron, and can then be cut with the file and forged, and
       in general can be made to assume any shape, like wrought iron. In
       this soft condition it is called _annealed steel_. The transition
       from tempered to annealed steel thus takes place in a similar way
       to the transition from white to grey cast iron. Steel, when
       homogeneous, has considerable lustre, and such a fine granular
       structure that it takes a very high polish. Its fracture clearly
       shows the granular nature of its structure. The possibility of
       tempering steel enables it to be used for making all kinds of
       cutting instruments, because annealed steel can be forged, turned,
       drawn (under rollers, for instance, for making rails, bars, &c.),
       filed, &c., and it may then be tempered, ground and polished. The
       method and temperature of tempering and annealing steel determine
       its hardness and other qualities. Steel is generally tempered to
       the required degree of hardness in the following manner: It is
       first strongly heated (for instance, up to 600°), and then plunged
       into water--that is, hardened by rapid cooling (it then becomes as
       brittle as glass). It is then heated until the surface assumes a
       definite colour, and finally cooled either quickly or slowly. When
       steel is heated up to 220°, its surface acquires a yellow colour
       (surgical instruments); it first of all becomes straw-coloured
       (razors, &c.), and then gold-coloured; then at a temperature of
       250° it becomes brown (scissors), then red, then light blue at
       285° (springs), then indigo at 300° (files), and finally sea-green
       at about 340°. These colours are only the tints of thin films,
       like the hues of soap bubbles, and appear on the steel because a
       thin layer of oxides is formed over its surface. Steel rusts more
       slowly than wrought iron, and is more soluble in acids than cast
       iron, but less so than wrought iron. Its specific gravity is about
       7·6 to 7·9.

       As regards the formation of steel, it was a long time before the
       process of cementation was thoroughly understood, because in this
       case infusible charcoal permeates unfused wrought iron. Caron
       showed that this permeation depends on the fact that the charcoal
       used in the process contains alkalis, which, in the presence of
       the nitrogen of the air, form metallic cyanides; these being
       volatile and fusible, permeate the iron, and, giving up their
       carbon to it, serve as the material for the formation of steel.
       This explanation is confirmed by the fact that charcoal without
       alkalis or without nitrogen will not cement iron. The charcoal
       used for cementation acts badly when used over again, as it has
       lost alkali. The very volatile ammonium cyanide easily conduces to
       the formation of steel. Although steel is also formed by the
       action of cyanogen compounds, nevertheless it does not contain
       more nitrogen than cast or wrought iron (0·01 p.c.), and these
       latter contain it because their ores contain titanium, which
       combines directly with nitrogen. Hence the part played by nitrogen
       in steel is but an insignificant one. It may be useful here to add
       some information taken from Caron's treatise concerning the
       influence of foreign matter on the quality of steel. The principal
       properties of steel are those of tempering and annealing. The
       compounds of iron with silicon and boron have not these
       properties. They are more stable than the carbon compound, and
       this latter is capable of changing its properties; because the
       carbon in it either enters into combination or else is disengaged,
       which determines the condition of hardness or softness of steel,
       as in white and grey cast iron. When slowly cooled, steel splits
       up into a mixture of soft and carburetted iron; but, nevertheless,
       the carbon does not separate from the iron. If such steel be again
       heated, it forms a uniform compound, and hardens when rapidly
       cooled. If the same steel as before be taken and heated a long
       time, then, after being slowly cooled, it becomes much more
       soluble in acid, and leaves a residue of pure carbon. This shows
       that the combination between the carbon and iron in steel becomes
       destroyed when subjected to heat, and the steel becomes iron mixed
       with carbon. Such _burnt_ steel cannot be tempered, but may be
       corrected by continued forging in a heated condition, which has
       the effect of redistributing the carbon equally throughout the
       whole mass. After the forging, if the iron is pure and the carbon
       has not been burnt out, steel is again formed, which may be
       tempered. If steel be repeatedly or strongly heated, it becomes
       burnt through and cannot be tempered or annealed; the carbon
       separates from the iron, and this is effected more easily if the
       steel contains other impurities which are capable of forming
       stable combinations with iron, such as silicon, sulphur, or
       phosphorus. If there be much silicon, it occupies the place of the
       carbon, and then continued forging will not induce the carbon once
       separated to re-enter into combination. Such steel is easily burnt
       through and cannot be corrected; when burnt through, it is hard
       and cannot be annealed--this is tough steel, an inferior kind.
       Iron which contains sulphur and phosphorus cements badly, combines
       but little with carbon, and steel of this kind is brittle, both
       hot and cold. Iron in combination with the above-mentioned
       substances cannot be annealed by slow cooling, showing that these
       compounds are more stable than those of carbon and iron, and
       therefore they prevent the formation of the latter. Such metals as
       tin and zinc combine with iron, but not with carbon, and form a
       brittle mass which cannot be annealed and is deleterious to steel.
       Manganese and tungsten, on the contrary, are capable of combining
       with charcoal; they do not hinder the formation of steel, but even
       remove the injurious effects of other admixtures (by transforming
       these admixed substances into new compounds and slags), and are
       therefore ranked with the substances which act beneficially on
       steel; but, nevertheless, the best steel, which is capable of
       renewing most often its primitive qualities after burning or hot
       forging, is the purest. The addition of Ni, Cr, W, and certain
       other metals to steel renders it very suitable for certain special
       purposes, and is therefore frequently made use of.

       It is worthy of attention that steel, besides temper, possesses
       many variable properties, a review of which may be made in the
       classification of the _sorts of steel_ (1878, Cockerell). (1)
       _Very mild steel_ contains from 0·05 to 0·20 p.c. of carbon,
       breaks with a weight of 40 to 50 kilos per square millimetre, and
       has an extension of 20 to 30 p.c.; it may be welded, like wrought
       iron, but cannot be tempered; is used in sheets for boilers,
       armour plate and bridges, nails, rivets, &c., as a substitute for
       wrought iron; (2) _mild steel_, from 0·20 to 0·35 p.c. of carbon,
       resistance to tension 50 to 60 kilos, extension 15 to 20 p.c., not
       easily welded, and tempers badly, used for axles, rails, and
       railway tyres, for cannons and guns, and for parts of machines
       destined to resist bending and torsion; (3) _hard steel_, carbon
       0·35 to 0·50 p.c., breaking weight 60 to 70 kilos per square
       millimetre, extension 10 to 15 p.c., cannot be welded, takes a
       temper; used for rails, all kinds of springs, swords, parts of
       machinery in motion subjected to friction, spindles of looms,
       hammers, spades, hoes, &c.; (4) _very hard steel_, carbon 0·5 to
       0·65 p.c., tensile breaking weight 70 to 80 kilos, extension 5 to
       10 p.c., does not weld, but tempers easily; used for small
       springs, saws, files, knives and similar instruments.

       The properties of ordinary _wrought iron_ are well known. The best
       iron is the most tenacious--that is to say, that which does not
       break up when struck with the hammer or bent, and yet at the same
       time is sufficiently hard. There is, however, a distinction
       between hard and soft iron. Generally the softest iron is the most
       tenacious, and can best be welded, drawn into wire, sheets, &c.
       Hard, especially tough, iron is often characterised by its
       breaking when bent, and is therefore very difficult to work, and
       objects made from it are less serviceable in many respects. Soft
       iron is most adapted for making wire and sheet iron and such small
       objects as nails. Soft iron is characterised by its attaining a
       fibrous fracture after forging, whilst tough iron preserves its
       granular structure after this operation. Certain sorts of iron,
       although fairly soft at the ordinary temperature, become brittle
       when heated and are difficult to weld. These sorts are less
       suitable for being worked up into small objects. The variety of
       the properties of iron depends on the impurities which it
       contains. In general, the iron used in the arts still contains
       carbon and always a certain quantity of silicon, manganese,
       sulphur, phosphorus, &c. A variety in the proportion of these
       component parts changes the quality of the iron. In addition to
       this the change which soft wrought iron, having a fibrous
       structure, undergoes when subjected to repeated blows and
       vibrations is considerable; it then becomes granular and brittle.
       This to a certain degree explains the want of stability of some
       iron objects--such as truck axles, which must be renewed after a
       certain term of service, otherwise they become brittle. It is
       evident that there are innumerable intermediate transitions from
       wrought iron to steel and cast iron.

       At the present day the greater part of the cast iron manufactured
       is converted into steel, generally cast steel (Bessemer's and
       Martin's). I may add the Urals, Donetz district, and other parts
       of Russia offer the greatest advantages for the development of an
       iron industry, because these localities not only contain vast
       supplies of excellent iron ore, but also coal, which is necessary
       for smelting it.

  [11] According to information supplied by A. T. Skinder's experiments
       at the Oboukoff Steel Works, 140 volumes of liquid molten steel
       give 128 volumes of solid metal. By means of a galvanic current of
       great intensity and dense charcoal as one electrode, and iron as
       the other, Bernadoss welded iron and fused holes through sheet
       iron. Soft wrought iron, like steel and soft malleable cast iron,
       may be melted in Siemens' regenerative furnaces, and in furnaces
       heated with naphtha.

  [11 bis] Gore (1869), Tait, Barret, Tchernoff, Osmond, and others
       observed that at a temperature approaching 600°--that is, between
       dark and bright red heat--all kinds of wrought iron undergo a
       peculiar change called _recalescence_, _i.e._ a spontaneous rise
       of temperature. If iron be considerably heated and allowed to
       cool, it may be observed that at this temperature the cooling
       stops--that is, latent heat is disengaged, corresponding with a
       change in condition. The specific heat, electrical conductivity,
       magnetic, and other properties then also change. In tempering, the
       temperature of recalescence must not be reached, and so also in
       annealing, &c. It is evident that a change of the internal
       condition is here encountered, exactly similar to the transition
       from a solid to a liquid, although there is no evident physical
       change. It is probable that attentive study would lead to the
       discovery of a similar change in other substances.

  [12] The particles of steel are linked together or connected more
       closely than those of the other metals; this is shown by the fact
       that it only breaks with a tensile strain of 50-80 kilos per sq.
       mm., whilst wrought iron only withstands about 30 kilos, cast iron
       10, copper 35, silver 23, platinum 30, wood 8. The elasticity of
       iron, steel, and other metals is expressed by the so-called
       _coefficient of elasticity_. Let a rod be taken whose length is L;
       if a weight, P, be hung from the extremity of it, it will lengthen
       to _l_. The less it lengthens under other equal conditions, the
       more elastic the material, if it resumes its original length when
       the weight is removed. It has been shown by experiment that the
       increase in length _l_, due to elasticity, is directly
       proportional to the length L and the weight P, and inversely
       proportional to the section, but changes with the material. The
       coefficient of elasticity expresses that weight (in kilos per sq.
       mm.) under which a rod having a square section taken as 1 (we take
       1 sq. mm.) acquires double the length by tension. Naturally in
       practice materials do not withstand such a lengthening, under a
       certain weight they attain a limit of elasticity, _i.e._ they
       stretch permanently (undergo deformation). Neglecting fractions
       (as the elasticity of metals varies not only with the temperature,
       but also with forging, purity, &c.), the coefficient of elasticity
       of steel and iron is 20,000, copper and brass 10,000, silver
       7,000, glass 6,000, lead 2,000, and wood 1,200.

_The chemical properties of iron_ have been already repeatedly
mentioned in preceding chapters. Iron _rusts_ in air at the ordinary
temperature--that is to say, it becomes covered with a layer of iron
oxides. Here, without doubt, the moisture of the air plays a part,
because in dry air iron does not oxidise at all, and also because, more
particularly, ammonia is always found in iron rust; the ammonia must
arise from the action of the hydrogen of the water, at the moment of its
separation, on the nitrogen of the air. Highly-polished steel does not
rust nearly so readily, but if moistened with water, it easily becomes
coated with rust. As rust depends on the access of moisture, iron may be
preserved from rust by coating it with substances which prevent the
moisture having access to it. Thus arises the practice of covering iron
objects with paraffin,[13] varnish, oil, paints, or enamelling it with a
glassy-looking flux possessing the same coefficient of expansion as iron,
or with a dense scoria (formed by the heat of superheated steam), or with
a compact coating of various metals. Wrought iron (both as sheet iron and
in other forms), cast iron, and steel are often coated with tin, copper,
lead, nickel, and similar metals, which prevent contact with the air.
These metals preserve iron very effectually from rust if they form a
completely compact surface, but in those places where the iron becomes
exposed, either accidentally or from wear, rust appears much more quickly
than on a uniform iron surface, because, towards these metals (and also
towards the rust), the iron will then behave as an electro-positive pole
in a galvanic couple, and hence will attract oxygen. A coating of zinc
does not produce this inconvenience, because iron is electro-negative
with reference to zinc, in consequence of which galvanised iron does not
easily rust, and even an iron boiler containing some lumps of zinc rusts
less than one without zinc.[14] Iron oxidises at a high temperature,
forming _iron scale_, Fe_{3}O_{4}, composed of ferrous and ferric oxides,
and, as has been seen, decomposes water and acids with the evolution of
hydrogen. It is also capable of decomposing salts and oxides of other
metals, which property is applied in the arts for the extraction of
copper, silver, lead, tin, &c. For this reason iron is soluble in the
solutions of many salts--for instance, in cupric sulphate, with
precipitation of copper and formation of ferrous sulphate.[15] When iron
_acts on acids_ it always _forms compounds_ FeX_{2}--that is,
corresponding to the suboxide FeO--and answering to magnesium
compounds--and hence two atoms of hydrogen are replaced by one atom of
iron. Strongly oxidising acids like nitric acid may transform the ferrous
salt which is forming into the higher degree of oxidation or ferric salt
(corresponding with the sesquioxide, Fe_{2}O_{3}), but this is a
secondary reaction. Iron, although easily soluble in dilute nitric acid,
loses this property when plunged into strong fuming nitric acid; after
this operation it even loses the property of solubility in other acids
until the external coating formed by the action of the strong nitric acid
is mechanically removed. This condition of iron is termed the passive
state. _The passive condition_ of iron depends on the formation, on its
surface, of a coating of oxide due to the iron being acted on by the
lower oxides of nitrogen contained in the fuming nitric acid.[16] Strong
nitric acid which does not contain these lower oxides, does not render
iron passive, but it is only necessary to add some alcohol or other
reducing agent which forms these lower oxides in the nitric acid, and the
iron will assume the passive state.

  [13] Paraffin is one of the best preservatives for iron against
       oxidation in the air. I found this by experiments about 1860, and
       immediately published the fact. This method is now very generally
       applied.

  [14] See Chapter XVIII., Note 34 bis. Based on the rapid oxidation of
       iron and its increase in volume in the presence of water and salts
       of ammonium, a packing is used for water mains and steam pipes
       which is tightly hammered into the socket joints. This packing
       consists of a mixture of iron filings and a small quantity of
       sal-ammoniac (and sulphur) moistened with water; after a certain
       lapse of time, especially after the pipes have been used, this
       mass swells to such an extent that it hermetically seals the
       joints of the pipes.

  [15] Here, however, a ferric salt may also be formed (when all the iron
       has dissolved and the cupric salt is still in excess), because the
       cupric salts are reduced by ferrous salts. Cast iron is also
       dissolved.

  [16] Powdery reduced iron is passive with regard to nitric acid of a
       specific gravity of 1·37, but when heated the acid acts on it.
       This passiveness disappears in the magnetic field. Saint-Edme
       attributes the passiveness of iron (and nickel) to the formation
       of nitride of iron on the surface of the metal, because he
       observed that when heated in dry hydrogen ammonia is evolved by
       passive iron.

       Remsen observed that if a strip of iron be immersed in acid and
       placed in the magnetic field, it is principally dissolved at its
       middle part--that is, the acid acts more feebly at the poles.
       According to Étard (1891) strong nitric acid dissolves iron in
       making it passive, although the action is a very slow one.

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 FeX_{2}, and
in the presence of air and oxidising agents they change by degrees into
ferric salts of the type FeX_{3}. 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, FeSO_{4},[17] a white precipitate of ferrous hydroxide,
FeH_{2}O_{2}, 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, BaCO_{3}, 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 Fe_{2}Cl_{6} + 3BaCO_{3} + 3H_{2}O
= Fe_{2}O_{3},3H_{2}O + 3BaCl_{2} + 3CO_{2}. 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, AuCl_{3}, 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, Fe_{2}O_{3}
(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, FeSO_{4}, there is one equivalent of iron to one
equivalent of sulphur (in the sulphuric radicle), but in the neutral
ferric salt, Fe_{2}(SO_{4})_{3}, there is one equivalent of iron to one
and a half of sulphur in the form of the elements of sulphuric acid.[19]

  [17] _Iron vitriol_ or _green vitriol_, sulphate of iron or ferrous
       sulphate, generally crystallises from solutions, like magnesium
       sulphate, with seven molecules of water, FeSO_{4},7H_{2}O. This
       salt is not only formed by the action of iron on sulphuric acid,
       but also by the action of moisture and air on iron pyrites,
       especially when previously roasted (FeS_{2} + O_{2} = FeS +
       SO_{2}), and in this condition it easily absorbs the oxygen of
       damp air (FeS + O_{4} = FeSO_{4}). Green vitriol is obtained in
       many processes as a by-product. Ferrous sulphate, like all the
       ferrous salts, has a pale greenish colour hardly perceptible in
       solution. If it be desired to preserve it without change--that is,
       so as not to contain ferric compounds--it is necessary to keep it
       hermetically sealed. This is best done by expelling the air by
       means of sulphurous anhydride or ether; sulphurous anhydride,
       SO_{2}, removes oxygen from ferric compounds, which might be
       formed, and is itself changed into sulphuric acid, and hence the
       oxidation of the ferrous compound does not take place in its
       presence. Unless these precautions are taken, green vitriol turns
       brown, partly changing into the ferric salt. When turned brown, it
       is not completely soluble in water, because during its oxidation a
       certain amount of free insoluble ferric oxide is formed: 6FeSO_{4}
       + O_{3} = 2Fe_{2}(SO_{4})_{3} + Fe_{2}O_{3}. In order to cleanse
       such mixed green vitriol from the oxide, it is necessary to add
       some sulphuric acid and iron and boil the mixture; the ferric salt
       is then transformed into the ferrous state: Fe_{2}(SO_{4})_{3} +
       Fe = 3FeSO_{4}.

       Green vitriol is used for the manufacture of Nordhausen sulphuric
       acid (Chapter XX.), for preparing ferric oxide, in many dye works
       (for preparing the indigo vats and reducing blue indigo to white),
       and in many other processes; it is also a very good disinfectant,
       and is the cheapest salt from which other compounds of iron may be
       obtained.

       The other ferrous salts (excepting the yellow prussiate, which
       will be mentioned later) are but little used, and it is therefore
       unnecessary to dwell upon them. We will only mention _ferrous
       chloride_, which, in the crystalline state, has the composition
       FeCl_{2},4H_{2}O. It is easily prepared; for instance, by the
       action of hydrochloric acid on iron, and in the anhydrous state by
       the action of hydrochloric acid gas on metallic iron at a red
       heat. The anhydrous ferrous chloride then volatilises in the form
       of colourless cubic crystals. Ferrous oxalate (or the double
       potassium salt) acts as a powerful reducing agent, and is
       frequently employed in photography (as a developer).

  [18] Ferrous sulphate, like magnesium sulphate, easily forms double
       salts--for instance, (NH_{4})_{2}SO_{4},FeSO_{4},6H_{2}O. This
       salt does not oxidise in air so readily as green vitriol, and is
       therefore used for standardising KMnO_{4}.

  [19] The transformation of ferrous oxide into ferric oxide is not
       completely effected in air, as then only a part of the suboxide is
       converted into ferric oxide. Under these circumstances the
       so-called magnetic oxide of iron is generally produced, which
       contains atomic quantities of the suboxide and oxide--namely,
       FeO,Fe_{2}O_{3} = Fe_{3}O_{4}. This substance, as already
       mentioned, is found in nature and in iron scale. It is also formed
       when most ferrous and ferric salts are heated in air; thus, for
       instance, when ferrous carbonate, FeCO_{3} (native or the
       precipitate given by soda in a solution of FeX_{2}), is heated it
       loses the elements of carbonic anhydride, and magnetic oxide
       remains. This oxide of iron is attracted by the magnet, and is on
       this account called magnetic oxide, although it does not always
       show magnetic properties. If magnetic oxide be dissolved in any
       acid--for instance, hydrochloric--which does not act as an
       oxidising agent, a ferrous salt is first formed and ferric oxide
       remains, which is also capable of passing into solution. The best
       way of preparing the hydrate of the magnetic oxide is by
       decomposing a mixture of ferrous and ferric salts with ammonia; it
       is, however, indispensable to pour this mixture into the ammonia,
       and not _vice versâ_, as in that case the ferrous oxide would at
       first be precipitated alone, and then the ferric oxide. The
       compound thus formed has a bright green colour, and when dried
       forms a black powder. Other combinations of ferrous with ferric
       oxide are known, as are also compounds of ferric oxide with other
       bases. Thus, for instance, compounds are known containing 4
       molecules of ferrous oxide to 1 of ferric oxide, and also 6 of
       ferrous to 1 of ferric oxide. These are also magnetic, and are
       formed by heating iron in air. The magnesium compound
       MgO,Fe_{2}O_{3} is prepared by passing gaseous hydrochloric acid
       over a heated mixture of magnesia and ferric oxide. Crystalline
       magnesium oxide is then formed, and black, shiny, octahedral
       crystals of the above-mentioned composition. This compound is
       analogous to the aluminates--for instance, to spinel. Bernheim
       (1888) and Rousseau (1891) obtained many similar compounds of
       ferric oxide, and their composition apparently corresponds to the
       hydrates (Note 22) known for the oxide.

The most simple oxidising agent for transforming ferrous into ferric
salts is chlorine in the presence of water--for instance, 2FeCl_{2} +
Cl_{2} = Fe_{2}Cl_{6}, or, generally speaking, 2FeO + Cl_{2} + H_{2}O =
Fe_{2}O_{3} + 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, HMnO_{4}, in the presence of
acids--for example, 10FeSO_{4} + 2KMnO_{4} + 8H_{2}SO_{4} =
5Fe_{2}(SO_{4})_{3} + 2MnSO_{4} + K_{2}SO_{4} + 8H_{2}O. 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, Fe_{2}Cl_{6} + H_{2}S =
2FeCl_{2} + 2HCl + S. Sodium thiosulphate acts in a similar way:
Fe_{2}Cl_{6} + Na_{2}S_{2}O_{3} + H_{2}O = 2FeCl_{2} + Na_{2}SO_{4} +
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, Fe_{2}Cl_{6} + Zn = 2FeCl_{2} + ZnCl_{2}. Thus _the
transition from ferrous salts to ferric salts and vice versâ is always
possible_.[21]

  [20] Copper and cuprous salts also reduce ferric oxide to ferrous
       oxide, and are themselves turned into cupric salts. The essence of
       the reactions is expressed by the following equations: Fe_{2}O_{3}
       + Cu_{2}O = 2FeO + 2CuO; Fe_{2}O_{3} + Cu = 2FeO + CuO. This fact
       is made use of in analysing copper compounds, the quantity of
       copper being ascertained by the amount of ferrous salt obtained.
       An excess of ferric salt is required to complete the reaction.
       Here we have an example of reverse reaction; the ferrous oxide or
       its salt in the presence of alkali transforms the cupric oxide
       into cuprous oxide and metallic copper, as observed by Lovel,
       Knopp, and others.

  [21] We will here mention the reactions by means of which it may be
       ascertained whether the ferrous compound has been entirely
       converted into a ferric compound or _vice versâ_. There are two
       substances which are best employed for this purpose: potassium
       ferricyanide, FeK_{3}C_{6}N_{6}, and potassium thiocyanate, KCNS.
       The first salt gives with ferrous salts a blue precipitate of an
       insoluble salt, having a composition Fe_{5}C_{12}N_{12}; but with
       ferric salts it does not form any precipitate, and only gives a
       brown colour, and therefore when transforming a ferrous salt into
       a ferric salt, the completion of the transformation may be
       detected by taking a drop of the liquid on paper or on a porcelain
       plate and adding a drop of the ferricyanide solution. If a blue
       precipitate be formed, then part of the ferrous salt still
       remains; if there is none, the transformation is complete. The
       thiocyanate does not give any marked coloration with ferrous
       salts; but with ferric salts in the most diluted state it forms a
       bright red soluble compound, and therefore when transforming a
       ferric salt into a ferrous salt we must proceed as before, testing
       a drop of the solution with thiocyanate, when the absence of a red
       colour will prove the total transformation of the ferric salt into
       the ferrous state, and if a red colour is apparent it shows that
       the transformation is not yet complete.

_Ferric oxide_, or _sesquioxide of iron_, Fe_{2}O_{3}, 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 2Fe_{2}O_{3}3H_{2}O,
or Fe_{4}H_{6}O_{9}. 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
Fe_{2}O_{3},H_{2}O. 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]

  [22] The two ferric hydroxides are not only characterised by the
       above-mentioned properties, but also by the fact that the first
       hydroxide forms immediately with potassium ferrocyanide,
       K_{4}FeC_{6}N_{6}, a blue colour depending on the formation of
       Prussian blue, whilst the second hydroxide does not give any
       reaction whatever with this salt. The first hydroxide is entirely
       soluble in nitric, hydrochloric, and all other acids; whilst the
       second sometimes (not always) forms a brick-coloured liquid, which
       appears turbid and does not give the reactions peculiar to the
       ferric salts (Péan de Saint-Gilles, Scheurer-Kestner). In addition
       to this, when the smallest quantity of an alkaline salt is added
       to this liquid, ferric oxide is precipitated. Thus a colloidal
       solution is formed (hydrosol), which is exactly similar to silica
       hydrosol (Chapter XVII.), according to which example the hydrosol
       of ferric oxide may be obtained.

       If ordinary ferric hydroxide be dissolved in acetic acid, a
       solution of the colour of red wine is obtained, which has all the
       reactions characteristic of ferric salts. But if this solution
       (formed in the cold) be heated to the boiling-point, its colour is
       very rapidly intensified, a smell of acetic acid becomes apparent,
       and the solution then contains a new variety of ferric oxide. If
       the boiling of the solution be continued, acetic acid is evolved,
       and the modified ferric oxide is precipitated. If the evaporation
       of the acetic acid be prevented (in a closed or sealed vessel),
       and the liquid be heated for some time, the whole of the ferric
       hydroxide then passes into the insoluble form, and if some
       alkaline salt be added (to the hydrosol formed), the whole of the
       ferric oxide is then precipitated in its insoluble form. This
       method may be applied for separating ferric oxide from solutions
       of its salts.

       All phenomena observed respecting ferric oxide (colloidal
       properties, various forms, formation of double basic salts)
       demonstrate that this substance, like silica, alumina, lead
       hydroxide, &c., is polymerised, that the composition is
       represented by (Fe_{2}O_{3})_{_n_}.

The normal salts of the composition Fe_{2}X_{6} or FeX_{3} correspond
with ferric oxide--for example, the exceedingly volatile _ferric
chloride_, Fe_{2}Cl_{6}, which is easily prepared in the anhydrous state
by the action of chlorine on heated iron.[23] Such also is the _normal
ferric nitrate_, Fe_{2}(NO_{3})_{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, Fe_{2}(NO_{3})_{6},9H_{2}O, 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 HNO_{3}) 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 R_{2}O_{3}) from many other bases (of the form RO) whose
nitrates are far more stable. The ferric salts, FeX_{3}, in passing into
ferrous salts, act as oxidising agents, as is seen from the fact that
they not only liberate S from SH_{2}, but also iodine from KI like many
oxidising agents.[25 bis]

  [23] The ferric compound which is most used in practice (for instance,
       in medicine, for cauterising, stopping bleeding, &c.--Oleum
       Martis) is _ferric chloride_, Fe_{2}Cl_{6}, easily obtained by
       dissolving the ordinary hydrated oxide of iron in hydrochloric
       acid. It is obtained in the anhydrous state by the action of
       chlorine on heated iron. The experiment is carried on in a
       porcelain tube, and a solid _volatile substance_ is then formed in
       the shape of brilliant violet scales which very readily absorb
       moisture from the air, and when heated with water decompose into
       crystalline ferric oxide and hydrochloric acid: Fe_{2}Cl_{6} +
       3H_{2}O = 6HCl + Fe_{2}O_{3}. Ferric chloride is so volatile that
       the density of its vapour may be determined. At 440° it is equal
       to 164·0 referred to hydrogen; the formula Fe_{2}Cl_{6}
       corresponds with a density of 162·5. An aqueous solution of this
       salt has a brown colour. On evaporating and cooling this solution,
       crystals separate containing 6 or 12 molecules of H_{2}O. Ferric
       chloride is not only soluble in water, but also in alcohol
       (similarly to magnesium chloride, &c.) and in ether. If the latter
       solutions are exposed to the rays of the sun they become
       colourless, and deposit ferrous chloride, FeCl_{2}, chlorine being
       disengaged. After a certain lapse of time, the aqueous solutions
       of ferric chloride decompose with precipitation of a basic salt,
       thus demonstrating the instability of ferric chloride, like the
       other salts of ferric oxide (Note 22). This salt is much more
       stable in the form of double salts, like all the ferric salts and
       also the salts of many other feeble bases. Potassium or ammonium
       chloride forms with it very beautiful red crystals of a double
       salt, having the composition Fe_{2}Cl_{6},4KCl,2H_{2}O. When a
       solution of this salt is evaporated it decomposes, with separation
       of potassium chloride.

       B. Roozeboom (1892) studied in detail (as for CaCl_{2}, Chapter
       XIV., Note 50) the separation of different hydrates from saturated
       solutions of Fe_{2}Cl_{6} at various concentrations and
       temperatures; he found that there are 4 crystallohydrates with 12,
       7, 5, and 4 molecules of water. An orange yellow only slightly
       hygroscopic hydrate, Fe_{2}Cl_{6},12H_{2}O, is most easily and
       usually obtained, which melts at 37°; its solubility at different
       temperatures is represented by the curve BCD in the accompanying
       figure, where the point B corresponds to the formation, at -55°,
       of a cryohydrate containing about Fe_{2}Cl_{6} + 36H_{2}O, the
       point C corresponds to the melting-point (+37°) of the hydrate
       Fe_{2}Cl_{6},12H_{2}O, and the curve CD to the fall in the
       temperature of crystallisation with an increase in the amount of
       salt, or decrease in the amount of water (in the figure the
       temperatures are taken along the axis of abscissæ, and the amount
       of _n_ in the formula _n_Fe_{2}Cl_{6} + 100H_{2}O along the axis
       of ordinates). When anhydrous Fe_{2}Cl_{6} is added to the above
       hydrate (12H_{2}O), or some of the water is evaporated from the
       latter, very hygroscopic crystals of Fe_{2}Cl_{6},5H_{2}O
       (Fritsche) are formed; they melt at 56°, their solubility is
       expressed by the curve HJ, which also presents a small branch at
       the end J. This again gives the fall in the temperature of
       crystallisation with an increase in the amount of Fe_{2}Cl_{6}.
       Besides these curves and the solubility of the anhydrous salt
       expressed by the line KL (up to 100°, beyond which chlorine is
       liberated), Roozeboom also gives the two curves, EFG and JK,
       corresponding to the crystallohydrates, Fe_{2}Cl_{6},7H_{2}O
       (melts at +32°·5, that is lower than any of the others) and
       Fe_{2}Cl_{6},4H_{2}O (melts at 73°·5), which he discovered by a
       systematic research on the solutions of ferric chloride. The curve
       AB represents the separation of ice from dilute solutions of the
       salt.

       [Illustration: FIG. 95.--Diagram of the solubility of
       Fe_{2}Cl_{6}.]

       The researches of the same Dutch chemist upon the conditions of
       the formation of crystals from the double salt
       (NH_{4}Cl)_{4}Fe_{2}Cl_{6},2H_{2}O are even more perfect. This
       salt was obtained in 1839 by Fritsche, and is easily formed from a
       strong solution of Fe_{2}Cl_{6} by adding sal-ammoniac, when it
       separates in crimson rhombic crystals, which, after dissolving in
       water, only deposit again on evaporation, together with the
       sal-ammoniac.

       Roozeboom (1892) found that when the solution contains _b_
       molecules of Fe_{2}Cl_{6}, and _a_ molecules of NH_{4}Cl, per 100
       molecules H_{2}O, then at 15° one of the following separations
       takes place: (1) crystals, Fe_{2}Cl_{6},12H_{2}O, when _a_ varies
       between 0 and 11, and _b_ between 4·65 and 4·8, or (2) a mixture
       of these crystals and the double salt, when _a_ = 1·36, and _b_ =
       4·47, or (3) the double salt, Fe_{2}Cl_{6},4NH_{4}Cl,2H_{2}O, when
       _a_ varies between 2 and 11·8, and _b_ between 3·1 and 4·56, or
       (4) a mixture of sal-ammoniac with the iron salt (it crystallises
       in separate cubes, Retgers, Lehmann), when _a_ varies between 7·7
       and 10·9, and _b_ is less than 3·38, or (5) sal-ammoniac, when _a_
       = 11·88. And as in the double salt, _a_ : _b_ :: 4 : 1 it is
       evident that the double salt only separates out when the ratio _a_
       : _b_ is less than 4 : 1 (_i.e._ when Fe_{2}Cl_{6} predominates).
       The above is seen more clearly in the accompanying figure, where
       _a_, or the number of molecules of NH_{4}Cl per 100H_{2}O, is
       taken along the axis of abscissæ, and _b_, or the number of
       molecules of Fe_{2}Cl_{6}, along the ordinates. The curves ABCD
       correspond to saturation and present an iso-therm of 15°. The
       portion AB corresponds to the separation of chloride of iron (the
       ascending nature of this curve shows that the solubility of
       Fe_{2}Cl_{6} is increased by the presence of NH_{4}Cl, while that
       of NH_{4}Cl decreases in the presence of Fe_{2}Cl_{6}), the
       portion BC to the double salt, and the portion CD to a mixture of
       sal-ammoniac and ferric chloride, while the straight line OF
       corresponds to the ratio Fe_{2}Cl_{6},4NH_{4}Cl, or _a_ : _b_ :: 4
       : 1. The portion CE shows that more double salt may be introduced
       into the solution without decomposition, but then the solution
       deposits a mixture of sal-ammoniac and ferric chloride (_see_
       Chapter XXIV. Note 9 ^{bis}). If there were more such
       well-investigated cases of solutions, our knowledge of double
       salts, solutions, the influence of water, equilibria, isomorphous
       mixtures, and such-like provinces of chemical relations might be
       considerably advanced.

       [Illustration: FIG. 96.--Diagram of the formation, at 15°, of the
       double salt Fe_{2}Cl_{6}4NH_{4}Cl_{2}H_{2}O or
       Fe(NH_{4})_{2}Cl_{5}H_{2}O. (After Roozeboom.)]

  [24] The normal ferric salts are decomposed by heat and even by water,
       forming basic salts, which may be prepared in various ways.
       Generally ferric hydroxide is dissolved in solutions of ferric
       nitrate; if it contains a double quantity of iron the basic salt
       is formed which contains Fe_{2}O_{3} (in the form of hydroxide) +
       2Fe_{2}(NO_{3})_{6} = 3Fe_{2}O(NO_{3})_{4}, a salt of the type
       Fe_{2}OX_{4}. Probably water enters into its composition. With
       considerable quantities of ferric oxide, insoluble basic salts are
       obtained containing various amounts of ferric hydroxide. Thus when
       a solution of the above-mentioned basic acid is boiled, a
       precipitate is formed containing
       4(Fe_{2}O_{3})_{8},2(N_{2}O_{5}),3H_{2}O, which probably contains
       2Fe_{2}O_{2}(NO_{3})_{2} + 2Fe_{2}O_{3},3H_{2}O. If a solution of
       basic nitrate be sealed in a tube and then immersed in boiling
       water, the colour of the solution changes just in the same way as
       if a solution of ferric acetate had been employed (Note 22). The
       solution obtained smells strongly of nitric acid, and on adding a
       drop of sulphuric or hydrochloric acid the insoluble variety of
       hydrated ferric oxide is precipitated.

       Normal ferric _orthophosphate_ is soluble in sulphuric,
       hydrochloric, and nitric acids, but insoluble in others, such as,
       for instance, acetic acid. The composition of this salt in the
       anhydrous state is FePO_{4}, because in orthophosphoric acid there
       are three atoms of hydrogen, and iron, in the ferric state,
       replaces the three atoms of hydrogen. This salt is obtained from
       ferric acetate, which, with disodium phosphate, forms a _white
       precipitate_ of FePO_{4}, containing water. If a solution of
       ferric chloride (yellowish-red colour) be mixed with a solution of
       sodium acetate in excess, the liquid assumes an intense brown
       colour which demonstrates the formation of a certain quantity of
       ferric acetate; then the disodium phosphate directly forms a white
       gelatinous precipitate of ferric phosphate. By this means the
       whole of the iron may be precipitated, and the liquid which was
       brown then becomes colourless. If this normal salt be dissolved in
       orthophosphoric acid, the crystalline acid salt
       FeH_{3}(PO_{4})_{2} is formed. If there be an excess of ferric
       oxide in the solution, the precipitate will consist of the basic
       salt. If ferric phosphate be dissolved in hydrochloric acid, and
       ammonia be added, a salt is precipitated on heating which, after
       continued washing in water and heating (to remove the water), has
       the composition Fe_{4}P_{2}O_{11}--that is,
       2Fe_{2}O_{3},P_{2}O_{5}. In an aqueous condition this salt may be
       considered as ferric hydroxide, Fe_{2}(OH)_{6}, in which (OH)_{3}
       is replaced by the equivalent group PO_{4}. Whenever ammonia is
       added to a solution containing an excess of ferric salt and a
       certain amount of phosphoric acid, a precipitate is formed
       containing the whole of the phosphoric acid in the mass of the
       ferric oxide.

       Ferric oxide is characterised as a feeble base, and also by the
       fact of its forming double salts--for instance, _potassium iron
       alum_, which has a composition
       Fe_{2}(SO_{4})_{3},K_{2}SO_{4},24H_{2}O or
       FeK(SO_{4})_{2},12H_{2}O. It is obtained in the form of almost
       colourless or light rose-coloured large octahedra of the regular
       system by simply mixing solutions of potassium sulphate and the
       ferric sulphate obtained by dissolving ferric oxide in sulphuric
       acid.

  [25] It would seem that all normal ferric salts are colourless, and
       that the brown colour which is peculiar to the solutions is really
       due to basic ferric salts. A remarkable example of the apparent
       change of colour of salts is represented by the ferrous and ferric
       oxalates. The former in a dry state has a yellow colour, although
       as a rule the ferrous salts are green, and the latter is
       colourless or pale green. When the normal ferric salt is dissolved
       in water it is, like many salts, probably decomposed by the water
       into acid and basic salts, and the latter communicates a brown
       colour to the solution. Iron alum is almost colourless, is easily
       decomposed by water, and is the best proof of our assertion. The
       study of the phenomena peculiar to ferric nitrate might, in my
       opinion, give a very useful addition to our knowledge of the
       aqueous solutions of salts in general.

  [25 bis] The reaction FeX_{3} + KI = FeX_{2} + KX + I proceeds
       comparatively slowly in solutions, is not complete (depends upon
       the mass), and is reversible. In this connection we may cite the
       following data from Seubert and Rohrer's (1894) comprehensive
       researches. The investigations were conducted with solutions
       containing 1/10 gram--equivalent weights of Fe_{2}(SO_{4})_{3}
       (_i.e._ containing 20 grams of salt per litre), and a
       corresponding solution of KI; the amount of iodine liberated being
       determined (after the addition of starch) by a solution (also 1/10
       normal) of Na_{2}S_{2}O_{3} (_see_ Chapter XX., Note 42). The
       progress of the reaction was expressed by the amount of liberated
       iodine in percentages of the theoretical amount. For instance, the
       following amount of iodide of potassium was decomposed when
       Fe_{2}(SO_{4})_{3} + 2_n_KI was taken:

               _n_ =        1     2     3     6    10    20
           After 15´      11·4  26·3  40·6  73·5  91·6  96·0
             "   30´      14·0  35·8  47·8  78·5  94·3  97·4
             "    1 hour  19·0  42·7  56·0  84·0  95·7  97·6
             "   10   "   32·6  56·0  75·7  93·2  96·5  97·6
             "   48   "   39·4  67·7  82·6  93·4  96·6  97·6

       Similar results were obtained for FeCl_{3}, but then the amount of
       iodine liberated was somewhat greater. Similar results were also
       obtained by increasing the mass of FeX_{3} per KI, and by
       replacing it by HI (_see_ Chapter XXI., Note 26).

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 K_{2}FeO_{4} is formed, and therefore the hydrate
corresponding with this salt should have the composition H_{2}FeO_{4}. It
is called _ferric acid_. Its anhydride ought to contain FeO_{3} or
Fe_{2}O_{6}--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 (2K_{2}FeO_{4} + 5H_{2}SO_{4} =
2K_{2}SO_{4} + Fe_{2}(SO_{4})_{3} + 5H_{2}O + O_{3}), 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:

  2CuSO_{4} + 2K_{2}FeO_{4} = 2K_{2}SO_{4} + O_{3} + Fe_{2}O_{3} + 2CuO.

Both these oxides are of course deposited in the form of hydrates. This
shows that not only the hydrate H_{2}FeO_{4}, but also the salts of the
heavy metals corresponding with this higher oxide of iron, are not formed
by reactions of double decomposition. The solution of potassium ferrate
naturally acts as a powerful oxidising agent; for instance, it transforms
manganous oxide into the dioxide, sulphurous into sulphuric acid, oxalic
acid into carbonic anhydride and water, &c.[26]

  [26] If chlorine be passed through a strong solution of potassium
       hydroxide in which hydrated ferric oxide is suspended, the turbid
       liquid acquires a dark pomegranate-red colour and contains
       potassium ferrate: 10KHO + Fe_{2}O_{3} + 3Cl_{2} = 2K_{2}FeO_{4} +
       6KCl + 5H_{2}O. The chlorine must not be in excess, otherwise the
       salt is again decomposed, although the mode of decomposition is
       unknown; however, ferric chloride and potassium chlorate are
       probably formed. Another way in which the above-described salt is
       formed is also remarkable; a galvanic current (from 6 Grove
       elements) is passed through cast-iron and platinum electrodes into
       a strong solution of potassium hydroxide. The cast-iron electrode
       is connected with the positive pole, and the platinum electrode is
       surrounded by a porous earthenware cylinder. Oxygen would be
       evolved at the cast-iron electrode, but it is used up in
       oxidation, and a dark solution of potassium ferrate is therefore
       formed about it. It is remarkable that the cast iron cannot be
       replaced by wrought iron.

Iron thus combines with oxygen in three proportions: RO, R_{2}O_{3},
and RO_{3}. It might have been expected that there would be intermediate
stages RO_{2} (corresponding to pyrites FeS_{2}) and R_{2}O_{5}, but for
iron these are unknown.[26 bis] The lower oxide has a distinctly basic
character, the higher is feebly acid. The only one which is stable in the
free state is ferric oxide, Fe_{2}O_{3}; the suboxide, FeO, absorbs
oxygen, and ferric anhydride, FeO_{3}, evolves it. It is also the same
for other elements; the character of each is determined by the relative
degree of stability of the known oxides. The salts FeX_{2} correspond
with the suboxide, the salts FeX_{3} or Fe_{2}X_{6} with the sesquioxide,
and FeX_{6} represents those of ferric acid, as its potassium salt is
FeO_{2}(OK)_{2}, corresponding with K_{2}SO_{4}, K_{2}MnO_{4},
K_{2}CrO_{4}, &c. Iron therefore forms compounds of the types FeX_{2},
FeX_{3}, and FeX_{6}, but this latter, like the type NX_{5}, does not
appear separately, but only when X represents heterogeneous elements or
groups; for instance, for nitrogen in the form of NO_{2}(OH), NH_{4}Cl,
&c., for iron in the form of FeO_{2}(OK)_{2}. But still the type FeX_{6}
exists, and therefore FeX_{2} and FeX_{3} are compounds which, like
ammonia, NH_{3}, are capable of further combinations up to FeX_{6}; this
is also seen in the property of ferrous and ferric salts of forming
compounds with water of crystallisation, besides double and basic salts,
whose stability is determined by the quality of the elements included in
the types FeX_{2} and FeX_{3}.[26 tri] It is therefore to be expected
that there should be complex compounds derived from ferrous and ferric
oxides. Amongst these the series of cyanogen compounds is particularly
interesting; their formation and character is not only determined by the
property which iron possesses of forming complex types, but also by the
similar faculty of the cyanogen compounds, which, like nitriles (Chapter
IX.), have clearly developed properties of polymerisation and in general
of forming complex compounds.[27]

  [26 bis] When Mond and his assistants obtained the remarkable volatile
       compound Ni(CO)_{4} (described later, Chapter XXII.), it was shown
       subsequently by Mond and Quincke (1891), and also by Berthelot,
       that iron, under certain conditions, in a stream of carbonic
       oxide, also volatilises and forms a compound like that given by
       nickel. Roscoe and Scudder then showed that when water gas is
       passed through and kept under pressure (8 atmospheres) in iron
       vessels a portion of the iron volatilises from the sides of the
       vessel, and that when the gas is burnt it deposits a certain
       amount of oxides of iron (the same result is obtained with
       ordinary coal gas which contains a small amount of CO). To obtain
       the _volatile compound of iron with carbonic oxide_, Mond prepared
       a finely divided iron by heating the oxalate in a stream of
       hydrogen, and after cooling it to 80°-45° he passed CO over the
       powder. The iron then formed (although very slowly) a volatile
       compound containing Fe(CO)_{5} (as though it answered to a very
       high type, FeX_{10}), which when cooled condenses into a liquid
       (slightly coloured, probably owing to incipient decomposition),
       sp. gr. 1·47, which solidifies at -21°, boils at about 103°, and
       has a vapour density (about 6·5 with respect to air) corresponding
       to the above formula; it decomposes at 180°. Water and dilute
       acids do not act upon it, but it decomposes under the action of
       light and forms a hard, non-volatile crystalline yellow compound
       Fe_{2}(CO)_{7} which decomposes at 80° and again forms Fe(CO)_{5}.

  [26 tri] When the molecular Fe_{2}Cl_{6} is produced instead of
       FeCl_{3} this complication of the type also occurs.

  [27] Some light may be thrown upon the faculty of Fe of forming various
       compounds with CN, by the fact that Fe not only combines with
       carbon but also with nitrogen. _Nitride of iron_ Fe_{2}N was
       obtained by Fowler by heating finely powdered iron in a stream of
       NH_{3} at the temperature of melting lead.

_In the cyanogen compounds of iron_, two degrees might be expected:
Fe(CN)_{2}, corresponding with ferrous oxide, and Fe(CN)_{3},
corresponding with ferric oxide. There are actually, however, many other
known compounds, intermediate and far more complex. They correspond with
the double salts so easily formed by metallic cyanides. The two following
double salts are particularly well known, very stable, often used, and
easily prepared. _Potassium ferrocyanide_ or _yellow prussiate of
potash_, a double salt of cyanide of potassium and ferrous cyanide, has
the composition FeC_{2}N_{2},4KCN; its crystals contain 3 mol. of water:
K_{4}FeC_{6}N_{6},3H_{2}O. The other is _potassium ferricyanide_ or _red
prussiate of potash_. It is also known as _Gmelin's salt_, and contains
cyanide of potassium with ferric cyanide; its composition is
Fe(CN)_{3},3KCN or K_{3}FeC_{6}N_{6}. Its crystals do not contain water.
It is obtained from the first by the action of chlorine, which removes
one atom of the potassium. A whole series of other ferrocyanic compounds
correspond with these ordinary salts.

Before treating of the preparation and properties of these two
remarkable and very stable salts, it must be observed that with ordinary
reagents neither of them gives the same double decompositions as the
other ferrous and ferric salts, and they both present a series of
remarkable properties. Thus these salts have a neutral reaction, are
unchanged by air, dilute acids, or water, unlike potassium cyanide and
even some of its double salts. When solutions of these salts are treated
with caustic alkalis, they do not give a precipitate of ferrous or ferric
hydroxides, neither are they precipitated by sodium carbonate. This led
the earlier investigators to recognise special independent groupings in
them. The yellow prussiate was considered to contain the complex radicle
FeC_{6}N_{6} combined with potassium, namely with K_{4}, and K_{3} was
attributed to the red prussiate. This was confirmed by the fact that
whilst in both salts any other metal, even hydrogen, might be substituted
for potassium, the iron remained unchangeable, just as nitrogen in
cyanogen, ammonium, and nitrates does not enter into double
decomposition, being in the state of the complex radicles CN, NH_{4},
NO_{2}. Such a representation is, however, completely superfluous for the
explanation of the peculiarities in the reactions of such compounds as
double salts. If a magnesium salt which can be precipitated by potassium
hydroxide does not form a precipitate in the presence of ammonium
chloride, it is very clear that it is owing to the formation of a soluble
double salt which is not decomposed by alkalis. And there is no necessity
to account for the peculiarity of reaction of a double salt by the
formation of a new complex radicle. In the same way also, in the presence
of an excess of tartaric acid, cupric salts do not form a precipitate
with potassium hydroxide, because a double salt is formed. These
peculiarities are more easily understood in the case of cyanogen
compounds than in all others, because all cyanogen compounds, as
unsaturated compounds, show a marked tendency to complexity. This
tendency is satisfied in double salts. The appearance of a peculiar
character in double cyanides is the more easily understood since in the
case of potassium cyanide itself, and also in hydrocyanic acid, a great
many peculiarities have been observed which are not encountered in those
haloid compounds, potassium chloride and hydrochloric acid, with which it
was usual to compare cyanogen compounds. These peculiarities become more
comprehensible on comparing cyanogen compounds with ammonium compounds.
Thus in the presence of ammonia the reactions of many compounds change
considerably. If in addition to this it is remembered that the presence
of many carbon (organic) compounds frequently completely disturbs the
reaction of salts, the peculiarities of certain double cyanides will
appear still less strange, because they contain carbon. The fact that the
presence of carbon or another element in the compound produces a change
in the reactions, may be compared to the action of oxygen, which, when
entering into a combination, also very materially changes the nature of
reactions. Chlorine is not detected by silver nitrate when it is in the
form of potassium chlorate, KClO_{3}, as it is detected in potassium
chloride, KCl. The iron in ferrous and ferric compounds varies in its
reactions. In addition to the above-mentioned facts, consideration ought
to be given to the circumstance that the easy mutability of nitric acid
undergoes modification in its alkali salts, and in general the properties
of a salt often differ much from those of the acid. Every double salt
ought to be regarded as a peculiar kind of saline compound: potassium
cyanide is, as it were, a basic, and ferrous cyanide an acid, element.
They may be unstable in the separate state, but form a stable double
compound when combined together; the act of combination disengages the
energy of the elements, and they, so to speak, saturate each other. Of
course, all this is not a definite explanation, but then the supposition
of a special complex radicle can even less be regarded as such.

Potassium ferrocyanide, K_{4}FeC_{6}N_{6}, is very easily formed by
mixing solutions of ferrous sulphate and potassium cyanide. First, a
white precipitate of ferrous cyanide, FeC_{2}N_{2}, is formed, which
becomes blue on exposure to air, but is soluble in an excess of potassium
cyanide, forming the ferrocyanide. The same yellow prussiate is obtained
on heating animal nitrogenous charcoal or animal matters--such as horn,
leather cuttings, &c.--with potassium carbonate in iron vessels,[27 bis]
the mass formed being afterwards boiled with water with exposure to air,
potassium cyanide first appearing, which gives yellow prussiate. The
animal charcoal may be exchanged for wood charcoal, permeated with
potassium carbonate and heated in nitrogen or ammonia; the mass thus
produced is then boiled in water with ferric oxide.[28] In this manner it
is manufactured on the large scale, and is called 'yellow prussiate'
('prussiate de potasse,' Blutlaugensalz).

  [27 bis] The sulphur of the animal refuse here forms the compound
       FeKS_{2}, which by the action of potassium cyanide yields
       potassium sulphide, thiocyanate, and ferrocyanide.

  [28] Potassium ferrocyanide may also be obtained from Prussian blue by
       boiling with a solution of potassium hydroxide, and from the
       ferricyanide by the action of alkalis and reducing substances
       (because the red prussiate is a product of oxidation produced by
       the action of chlorine: a ferric salt is reduced to a ferrous
       salt), &c. In many works (especially in Germany and France) yellow
       prussiate is prepared from the mass, containing oxide of iron, and
       employed for purifying coal gas (Vol. I., p. 361), which generally
       contains cyanogen compounds. About 2 p.c. of the nitrogen
       contained in coal is converted into cyanogen, which forms Prussian
       blue and thiocyanates in the mass used for purifying the gas. On
       evaporation the solution yields large yellow crystals containing 3
       molecules of water, which is easily expelled by heating above
       100°. 100 parts of water at the ordinary temperature are capable
       of dissolving 25 parts of this salt; its sp. gr. is 1·83. When
       ignited it forms potassium cyanide and iron carbide, FeC_{2}
       (Chapter XIII., Note 12). Oxidising substances change it into
       potassium ferricyanide. With strong sulphuric acid it gives
       carbonic oxide, and with dilute sulphuric acid, when heated,
       prussic acid is evolved according to the equation:
       2K_{4}FeC_{6}N_{6} + 3H_{2}SO_{4} = K_{2}Fe_{2}C_{6}N_{6} +
       3K_{2}SO_{4} + 6HCN; hence in the yellow prussiate K_{2} replaces
       Fe.

It is easy to substitute other metals for the potassium in the yellow
prussiate. The hydrogen salt or hydroferrocyanic acid, H_{4}FeC_{6}N_{6},
is obtained by mixing strong solutions of yellow prussiate and
hydrochloric acid. If ether be added and the air excluded, the acid is
obtained directly in the form of a white scarcely crystalline precipitate
which becomes blue on exposure to air (as ferrous cyanide does from the
formation of blue compounds of ferrous and ferric cyanides, and it is on
this account used in cotton printing). It is soluble in water and
alcohol, but not in ether, has marked acid properties, and decomposes
carbonates, which renders it easily possible to prepare ferrocyanides of
the metals of the alkalis and alkaline earths; these are readily soluble,
have a neutral reaction, and resemble the yellow prussiate. Solutions of
these salts form precipitates with the salts of other metals, because the
ferrocyanides of the heavy metals are insoluble. Here either the whole of
the potassium of the yellow prussiate, or only a part of it, is exchanged
for an equivalent quantity of the heavy metal. Thus, when a cupric salt
is added to a solution of yellow prussiate, a red precipitate is obtained
which still contains half the potassium of the yellow prussiate:

    K_{4}FeC_{6}N_{6} + CuSO_{4} = K_{2}CuFeC_{6}N_{6} + K_{2}SO_{4}.

But if the process be reversed (the salt of copper being then in excess)
the whole of the potassium will be exchanged for copper, forming a
reddish-brown precipitate, Cu_{2}FeC_{6}N_{6},9H_{2}O. This reaction and
those similar to it are very sensitive and may be used for detecting
metals in solution, more especially as the colour of the precipitate very
often shows a marked difference when one metal is exchanged for another.
Zinc, cadmium, lead, antimony, tin, silver, cuprous and aurous salts form
_white_ precipitates; cupric, uranium, titanium and molybdenum salts
_reddish-brown_; those of nickel, cobalt, and chromium, _green_
precipitates; _with ferrous salts_, ferrocyanide forms, as has been
already mentioned, a _white_ precipitate--namely, Fe_{2}FeC_{6}N_{6}, or
FeC_{2}N_{2}--which turns blue on exposure to air, and with ferric salts
a _blue precipitate_ called _Prussian blue_. Here the potassium is
replaced by iron, the reaction being expressed thus: 2Fe_{2}Cl_{6} +
3K_{4}FeC_{6}N_{6} = 12KCl + Fe_{4}Fe_{3}C_{18}N_{18}, the latter formula
expressing the composition of Prussian blue. It is therefore the compound
4Fe(CN)_{3} + 3Fe(CN)_{2}. The yellow prussiate is prepared in chemical
works on a large scale especially for the manufacture of this blue
pigment, which is used for dyeing cloth and other fabrics and also as one
of the ordinary blue paints. It is insoluble in water, and the stuffs are
therefore dyed by first soaking them in a solution of a ferric salt and
then in a solution of yellow prussiate. If however an excess of yellow
prussiate be present complete substitution between potassium and iron
does not occur, and _soluble Prussian blue_ is formed; KFe_{2}(CN)_{6} =
KCN,Fe(CN)_{2},Fe(CN)_{3}. This blue salt is colloidal, is soluble in
pure water, but insoluble and precipitated when other salts--for
instance, potassium or sodium chloride--are present even in small
quantities, and is therefore first obtained as a precipitate.[29]

  [29] Skraup obtained this salt both from potassium ferrocyanide with
       ferric chloride and from ferricyanide with ferrous chloride, which
       evidently shows that it contains iron in both the ferric and
       ferrous states. With ferrous chloride it forms Prussian blue, and
       with ferric chloride Turnbull's blue.

       Prussian blue was discovered in the beginning of the last century
       by a Berlin manufacturer, Diesbach. It was then prepared, as it
       sometimes is also at present, directly from potassium cyanide
       obtained by heating animal charcoal with potassium carbonate. The
       mass thus obtained is dissolved in water, alum is added to the
       solution in order to saturate the free alkali, and then a solution
       of green vitriol is added which has previously been sufficiently
       exposed to the air to contain both ferric and ferrous salts. If
       the solution of potassium cyanide be mixed with a solution
       containing both salts, Prussian blue will be formed, because it is
       a compound of ferrous cyanide, FeC_{2}N_{2}, and ferric cyanide,
       Fe_{2}C_{6}N_{6}. A ferric salt with potassium ferrocyanide forms
       a blue colour, because ferrous cyanide is obtained from the first
       salt and ferric cyanide from the second. During the preparation of
       this compound alkali must be avoided, as otherwise the precipitate
       would contain oxides of iron. Prussian blue has not a crystalline
       structure; it forms a blue mass with a copper-red metallic lustre.
       Both acids and alkalis act on it. The action is at first confined
       to the ferric salt it contains. Thus alkalis form ferric oxide and
       ferrocyanide in solution: 2Fe_{2}C_{6}N_{6},3FeC_{2}N_{2} + 12KHO
       = 2(Fe_{2}O_{3},3H_{2}O) + 3K_{4}FeC_{6}N_{6}. Various
       ferrocyanides may thus be prepared. Prussian blue is soluble in an
       aqueous solution of oxalic acid, forming blue ink. In air, when
       exposed to the action of light, it fades; but in the dark again
       absorbs oxygen and becomes blue, which fact is also sometimes
       noticed in blue cloth. An excess of potassium ferrocyanide renders
       Prussian blue soluble in water, although insoluble in various
       saline solutions--that is, it converts it into the soluble
       variety. Strong hydrochloric acid also dissolves Prussian blue.

Potassium ferricyanide, or _red prussiate_ of potash, K_{3}FeC_{6}N_{6},
is called 'Gmelin's salt,' because this savant obtained it by the action
of chlorine on a solution of the yellow prussiate: K_{4}FeC_{6}N_{6} + Cl
= K_{3}FeC_{6}N_{6} + KCl. The reaction is due to the ferrous salt being
changed by the action of the chlorine into a ferric salt. It separates
from solutions in anhydrous, well-formed prisms of a red colour, but the
solution has an olive colour; 100 parts of water, at 10°, dissolve 37
parts of the salt, and at 100°, 78 parts.[30] The red prussiate gives a
blue precipitate with ferrous salts, called _Turnbull's blue_, very much
like Prussian blue (and the soluble variety), because it also contains
ferrous cyanide and ferric cyanide, although in another proportion, being
formed according to the equation: 3FeCl_{2} + 2K_{3}FeC_{6}N_{6} = 6KCl +
Fe_{3}Fe_{2}C_{12}N_{12}, or 3FeC_{2}N_{2},Fe_{2}C_{6}N_{6}; in Prussian
blue we have Fe_{7}Cy_{18}, and here Fe_{5}Cy_{12}. A ferric salt ought
to form ferric cyanide Fe_{2}C_{6}N_{6}, with red prussiate, but ferric
cyanide is soluble, and therefore no precipitate is obtained, and the
liquid only becomes brown.[31]

  [30] An excess of chlorine must not be employed in preparing this
       compound, otherwise the reaction goes further. It is easy to find
       out when the action of the chlorine on potassium ferrocyanide must
       cease; it is only necessary to take a sample of the liquid and add
       a solution of a ferric salt to it. If a precipitate of Prussian
       blue is formed, more chlorine must be added, as there is still
       some undecomposed ferrocyanide, for the ferricyanide does not give
       a precipitate with ferric salts. Potassium ferricyanide, like the
       ferrocyanide, easily exchanges its potassium for hydrogen and
       various metals by double decomposition. With the salts of tin,
       silver, and mercury it forms yellow precipitates, and with those
       of uranium, nickel, cobalt, copper, and bismuth brown
       precipitates. The lead salt under the action of sulphuretted
       hydrogen forms lead sulphide and a hydrogen salt or acid,
       H_{3}FeC_{6}N_{6}, corresponding with potassium ferricyanide,
       which is soluble, crystallises in red needles, and resembles
       hydroferrocyanic acid, H_{4}FeC_{6}N_{6}. Under the action of
       reducing agents--for instance, sulphuretted hydrogen,
       copper--potassium ferricyanide is changed into ferrocyanide,
       especially in the presence of alkalis, and thus forms a rather
       energetic _oxidising agent_--capable, for instance, of changing
       manganous oxide into dioxide, bleaching tissues, &c.

  [31] It is important to mention a series of readily crystallisable
       salts formed by the action of nitric acid on potassium and other
       ferrocyanides and ferricyanides. These salt contain the elements
       of nitric oxide, and are therefore called _nitro-(nitroso)
       ferricyanides_ (_nitroprussides_). Generally a crystalline sodium
       salt is obtained, Na_{2}FeC_{5}N_{6}O,2H_{2}O. In its composition
       this salt differs from the red sodium salt, Na_{3}FeC_{6}N_{6}, by
       the fact that in it one molecule of sodium cyanide, NaCN, is
       replaced by nitric oxide, NO. In order to prepare it, potassium
       ferrocyanide in powder is mixed with five-sevenths of its weight
       of nitric acid diluted with an equal volume of water. The mixture
       is at first left at the ordinary temperature, and then heated on a
       water-bath. Here ferricyanide is first of all formed (as shown by
       the liquid giving a precipitate with ferrous chloride), which then
       disappears (no precipitate with ferrous chloride), and forms a
       green precipitate. The liquid, when cooled, deposits crystals of
       nitre. The liquid is then strained off and mixed with sodium
       carbonate, boiled, filtered, and evaporated; sodium nitrate and
       the salt described are deposited in crystals. It separates in
       prisms of a red colour. Alkalis and salts of the alkaline earths
       do not give precipitates: they are soluble, but the salts of iron,
       zinc, copper, and silver form precipitates where sodium is
       exchanged with these metals. It is remarkable that the sulphides
       of the alkali metals give with this salt an intense bright purple
       coloration. This series of compounds was discovered by Gmelin and
       studied by Playfair and others (1849).

       This series to a certain extent resembles the nitro-sulphide
       series described by Roussin. Here the primary compound consists of
       black crystals, which are obtained as follows:--Solutions of
       potassium hydrosulphide and nitrate are mixed, and the mixture is
       agitated whilst ferric chloride is added, then boiled and
       filtered; on cooling, _black crystals_ are deposited, having the
       composition Fe_{6}S_{3}(NO)_{10},H_{2}O (Rosenberg), or, according
       to Demel, FeNO_{2},NH_{2}S. They have a slightly metallic lustre,
       and are soluble in water, alcohol, and ether. They absorb the
       latter as easily as calcium chloride absorbs water. In the
       presence of alkalis these crystals remain unchanged, but with
       acids they evolve nitric oxides. There are several compounds which
       are capable of interchanging, and correspond with Roussin's salt.
       Here we enter into the series of the nitrogen compounds which have
       been as yet but little investigated, and will most probably in
       time form most instructive material for studying the nature of
       that element. These series of compounds are as unlike the usual
       saline compounds of inorganic chemistry as are organic
       hydrocarbons. There is no necessity to describe these series in
       detail, because their connection with other compounds is not yet
       clear, and they have not yet any application.

If chlorine and sodium are representatives of independent groups of
elements, the same may also be said of iron. Its nearest analogues show,
besides a similarity in character, a likeness as regards physical
properties and a proximity in atomic weight. Iron occupies a medium
position amongst its nearest analogues, both with respect to properties
and faculty of forming saline oxides, and also as regards atomic weight.
On the one hand, cobalt, 58, and nickel, 59, approach iron, 56; they are
metals of a more basic character, they do not form stable acids or higher
degrees of oxidation, and are a transition to copper, 63, and zinc, 65.
On the other hand, manganese, 55, and chromium, 52, are the nearest to
iron; they form both basic and acid oxides, and are a transition to the
metals possessing acid properties. In addition to having atomic weights
approximately alike, chromium, manganese, iron, cobalt, nickel, and
copper have also nearly the same specific gravity, so that the atomic
volumes and the molecules of their analogous compounds are also near to
one another (see table at the beginning of this volume). Besides this,
the likeness between the above-mentioned elements is also seen from the
following:

They form suboxides, RO, fairly energetic bases, isomorphous with
magnesia--for instance, the salt RSO_{4},7H_{2}O, akin to
MgSO_{4},7H_{2}O, and FeSO_{4},7H_{2}O, or to sulphates containing less
water; with alkali sulphates all form double salts crystallising with
6H_{2}O; all are capable of forming ammonium salts, &c. The lower oxides,
in the cases of nickel and cobalt, are tolerably stable, are not easily
oxidised (the nickel compound with more difficulty than cobalt, a
transition to copper); with manganese, and especially with chromium, they
are more easily oxidised than with iron and pass into higher oxides. They
also form oxides of the form R_{2}O_{3}, and with nickel, cobalt, and
manganese this oxide is very unstable, and is more easily reduced than
ferric oxide; but, in the case of chromium, it is very stable, and forms
the ordinary kind of salts. It is isomorphous with ferric oxide, forms
alums, is a feeble base, &c. Chromium, manganese, and iron are oxidised
by alkali and oxidising agents, forming salts like Na_{2}SO_{4}; but
cobalt and nickel are difficult to oxidise; their acids are not known
with any certainty, and are, in all probability, still less stable than
the ferrates. Cr, Mn and Fe form compounds R_{2}Cl_{6} which are like
Fe_{2}Cl_{6} in many respects; in Co this faculty is weaker and in Ni it
has almost disappeared. The cyanogen compounds, especially of manganese
and cobalt, are very near akin to the corresponding ferrocyanides. The
oxides of nickel and cobalt are more easily reduced to metal than those
of iron, but those of manganese and chromium are not reduced so easily as
iron, and the metals themselves are not easily obtained in a pure state;
they are capable of forming varieties resembling cast iron. The metals
Cr, Mn, Fe, Co, and Ni have a grey iron colour and are very difficult to
melt, but nickel and cobalt can be melted in the reverberatory furnace
and are more fusible than iron, whilst chromium is more difficult to melt
than platinum (Deville). These metals decompose water, but with greater
difficulty as the atomic weight rises, forming a transition to copper,
which does not decompose water. All the compounds of these metals have
various colours, which are sometimes very bright, especially in the
higher stages of oxidation.

These metals of the iron group are often met with together in nature.
Manganese nearly everywhere accompanies iron, and iron is always an
ingredient in the ores of manganese. Chromium is found principally as
chrome ironstone--that is, a peculiar kind of magnetic oxide in which
Fe_{2}O_{3} is replaced by Cr_{2}O_{3}.

Nickel and cobalt are as inseparable companions as iron and manganese.
The similarity between them even extends to such remote properties as
magnetic qualities. In this series of metals we find those which are the
most magnetic: iron, cobalt, and nickel. There is even a magnetic oxide
among the chromium compounds, such being unknown in the other series.
Nickel easily becomes passive in strong nitric acid. It absorbs hydrogen
in just the same way as iron. In short, in the series Cr, Mn, Fe, Co, and
Ni, there are many points in common although there are many differences,
as will be seen still more clearly on becoming acquainted with cobalt and
nickel.

In nature _cobalt_ is principally found in combination with arsenic
and sulphur. _Cobalt arsenide_, or _cobalt speiss_, CoAs_{2}, is found in
brilliant crystals of the regular system, principally in Saxony. _Cobalt
glance_, CoAs_{2}CoS_{2}, resembles it very much, and also belongs to the
regular system; it is found in Sweden, Norway, and the Caucasus.
_Kupfernickel_ is a nickel ore in combination with arsenic, but of a
different composition from cobalt arsenide, having the formula NiAs; it
is found in Bohemia and Saxony. It has a copper-red colour and is rarely
crystalline; it is so called because the miners of Saxony first mistook
it for an ore of copper (_Kupfer_), but were unable to extract copper
from it. _Nickel glance_, NiS_{2},NiAs_{2}, corresponding with cobalt
glance, is also known. Nickel accompanies the ores of cobalt and cobalt
those of nickel, so that both metals are found together. The ores of
cobalt are worked in the Caucasus in the Government of Elizavetopolsk.
Nickel ores containing aqueous hydrated nickel silicate are found in the
Ural (Revdansk). Large quantities of a similar ore are exported into
Europe from New Caledonia. Both ores contain about 12 per cent. Ni.
_Garnierite_, (RO)_{5}(SiO_{2})_{4}1-1/2H_{2}O, where R = Ni and Mg,
predominates in the New Caledonian ore. Large deposits of nickel have
been discovered in Canada, where the ore (as nickelous pyrites) is free
from arsenic. Cobalt is principally worked up into cobalt compounds, but
nickel is generally reduced to the metallic state, in which it is now
often used for alloys--for instance, for coinage in many European States,
and for plating other metals, because it does not oxidise. Cobalt
arsenide and cobalt glance are principally used for the preparation of
cobalt compounds; they are first sorted by discarding the rocky matter,
and then roasted. During this process most of the sulphur and arsenic
disappears; the arsenious anhydride volatilises with the sulphurous
anhydride and the metal also oxidises.[32] It is a simple matter to
obtain nickel and cobalt from their oxides. In order to obtain the
latter, solutions of their salts are treated with sodium carbonate and
the precipitated carbonates are heated; the suboxides are thus obtained,
and these latter are reduced in a stream of hydrogen, or even by heating
with ammonium chloride. They easily oxidise when in the state of powder.
When the chlorides of nickel and cobalt are heated in a stream of
hydrogen, the metal is deposited in brilliant scales. _Nickel is always
much more easily and quickly reduced than cobalt._ Nickel melts more
easily than cobalt, and this even furnishes a means of testing the
heating powers of a reverberatory furnace. Cobalt fuses at a temperature
only a little lower than that at which iron does. In general, cobalt is
nearer to iron than nickel, nickel being nearer to copper.[32 bis] Both
nickel and cobalt have magnetic properties like iron, but Co is less
magnetic than Fe, and Ni still less so. The specific gravity of nickel
reduced by hydrogen is 9·1 and that of cobalt 8·9. Fused cobalt has a
specific gravity of 8·5, the density of ordinary nickel being almost the
same. Nickel has a greyish silvery-white colour; it is brilliant and very
ductile, so that the finest wire may be easily drawn from it. This wire
has a resistance to tension equal to iron wire. The beautiful colour of
nickel, and the high polish which it is capable of receiving and
retaining, as it does not oxidise, render it a useful metal for many
purposes, and in many ways it resembles silver.[32 tri] It is now very
common to cover other metals with a layer of nickel (nickel plating).
This is done by a process of electro-plating, using a solution of a
nickel salt. The colour of cobalt is dark and redder; it is also ductile,
and has a greater tensile resistance than iron. Dilute acids act very
slowly on nickel and cobalt; nitric acid may be considered as the best
solvent for them. The solutions in every case contain salts corresponding
with the ferrous salts--that is, the _salts_ CoX_{2}, NiX_{2},
_correspond with the suboxides_ of these metals. These salts in their
types are similar to the magnesium salts. The salts of nickel when
crystallising with water have a green colour, and form bright green
solutions, but in the anhydrous state they most frequently have a yellow
colour. The salts of cobalt are generally rose-coloured, and generally
blue when in the anhydrous state. Their aqueous solutions are
rose-coloured. Cobaltous chloride is easily soluble in alcohol, and forms
a solution of an intense blue colour.[33]

  [32] The residue from the roasting of cobalt ores is called _zafflor_,
       and is often met with in commerce. From this the purer compounds
       of cobalt may be prepared. The ores of nickel are also first
       roasted, and the oxides dissolved in acid, nickelous salts being
       then obtained.

       The further treatment of cobalt and nickel ores is facilitated if
       the arsenic can be almost entirely removed, which may be effected
       by roasting the ore a second time with a small addition of nitre
       and sodium carbonate; the nitre combines with the arsenic, forming
       an arsenious salt, which may be extracted with water. The
       remaining mass is dissolved in hydrochloric acid, mixed with a
       small quantity of nitric acid. Copper, iron, manganese, nickel,
       cobalt, &c., pass into solution. By passing hydrogen sulphide
       through the solution, copper, bismuth, lead, and arsenic are
       deposited as metallic sulphides; but iron, cobalt, nickel, and
       manganese remain in solution. If an alkaline solution of bleaching
       powder be then added to the remaining solution, the whole of the
       manganese will first be deposited in the form of dioxide, then the
       cobalt as hydrated cobaltic oxide, and finally the nickel also. It
       is, however, impossible to rely on this method for effecting a
       complete separation, the more so since the higher oxides of the
       three above-mentioned metals have all a black colour; but, after a
       few trials, it will be easy to find how much bleaching powder is
       required to precipitate the manganese, and the amount which will
       precipitate all the cobalt. The manganese may also be separated
       from cobalt by precipitation from a mixture of the solutions of
       both metals (in the form of the 'ous' salts) with ammonium
       sulphide, and then treating the precipitate with acetic acid or
       dilute hydrochloric acid, in which manganese sulphide is easily
       soluble and cobalt sulphide almost insoluble. Further particulars
       relating to the separation of cobalt from nickel may be found in
       treatises on analytical chemistry. In practice it is usual to rely
       on the rough method of separation founded on the fact that nickel
       is more easily reduced and more difficult to oxidise than cobalt.
       The New Caledonian ore is smelted with CaSO_{4} and CaCO_{3} on
       coke, and a metallic regulus is obtained containing all the Ni,
       Fe, and S. This is roasted with SiO_{2}, which converts all the
       iron into slag, whilst the Ni remains combined with the S; this
       residue on further roasting gives NiO, which is reduced by the
       carbon to metallic Ni. The Canadian ore (a pyrites containing 11
       p.c. Ni) is frequently treated in America (after a preliminary
       dressing) by smelting it with Na_{2}SO_{4} and charcoal; the
       resultant fusible Na_{2}S then dissolves the CuS and FeS_{2},
       while the NiS is obtained in a bottom layer (Bartlett and
       Thomson's process) from which Ni is obtained in the manner
       described above.

       For manufacturing purposes somewhat impure cobalt compounds are
       frequently used, which are converted into _smalt_. This is glass
       containing a certain amount of cobalt oxide; the glass acquires a
       bright blue colour from this addition, so that when powdered it
       may be used as a blue pigment; it is also unaltered at high
       temperatures, so that it used to take the place now occupied by
       Prussian blue, ultramarine, &c. At present smalt is almost
       exclusively used for colouring glass and china. To prepare smalt,
       ordinary impure cobalt ore (zaffre) is fused in a crucible with
       quartz and potassium carbonate. A fused mass of cobalt glass is
       thus formed, containing silica, cobalt oxide, and potassium oxide,
       and a metallic mass remains at the bottom of the crucible,
       containing almost all the other metals, arsenic, nickel, copper,
       silver, &c. This metallic mass is called _speiss_, and is used as
       nickel ore for the extraction of nickel. Smalt usually contains 70
       p.c. of silica, 20 p.c. of potash and soda, and about 5 to 6 p.c.
       of cobaltous oxide; the remainder consisting of other metallic
       oxides.

  [32 bis] All we know respecting the relations of Co and Ni to Fe and Cu
       confirms the fact that Co is more closely related to Fe and Ni to
       Cu; and as the atomic weight of Fe = 56 and of Cu = 63, then
       according to the principles of the periodic system it would be
       expected that the atomic weight of Co would be about 59-60, whilst
       that of Ni should be greater than that of Co but less than that of
       Cu, _i.e._ about 50·5-60·5. However, as yet the majority of the
       determinations of the atomic weights of Co and Ni give a different
       result and show that a lower atomic weight is obtained for Ni than
       for Co. Thus K. Winkler (1894) obtained (employing metals
       deposited electrolytically and determining the amount of iodine
       which combined with them) Ni = 58·72 and Co = 59·37 (if H = 1 and
       I = 126·53). In my opinion this should not be regarded as proving
       that the principles of the periodic system cannot be applied in
       this instance, nor as a reason for altering the position of these
       elements in the system (_i.e._ by placing Ni after Fe, and Co next
       to Cu), because in the first place the figures given by different
       chemists (for instance, Zimmermann, Krüss, and others) are
       somewhat divergent, and in the second place the majority of the
       latest modes of determining the atomic weights of Co and Ni aim at
       finding what weights of these metals react with known weights of
       other elements without taking into account the faculty they have
       of absorbing hydrogen; since this faculty is more developed in Ni
       than in Co the hydrogen (occluded in Ni) should lower the atomic
       weight of Ni more than that of Co. On the whole, the question of
       the atomic weights of Co and Ni cannot yet be considered as
       decided, notwithstanding the numerous researches which have been
       made; still there can be no doubt that the atomic weights of these
       two metals are very nearly equal, and greater than that of Fe, but
       less than that of Cu. This question is of great interest, not only
       for completing our knowledge of these metals, but also for
       perfecting our knowledge of the periodic system of the elements.

  [32 tri] For instance, the alkalis may be fused in nickel vessels as
       well as in silver, because they have no action upon either metal.
       Nickel, like silver, is not acted upon by dilute acids. Only
       nitric acid dissolves both metals well. Nickel is harder, and
       fuses at a higher temperature than silver. For castings, a small
       quantity of magnesium (0·001 part by weight) is added to nickel to
       render it more homogeneous (just as aluminium is added to steel).
       Nickel forms many valuable alloys. Steel containing 3 p.c. Ni is
       particularly valuable, its limit of elasticity is higher and its
       hardness is greater; it is used for armour plate and other large
       pieces. The alloys of nickel, especially with copper and zinc
       (melchior, _see_ later), aluminium and silver, although used in
       certain cases, are now replaced by nickel-plated or
       nickel-deposited goods (deposited by electricity from a solution
       of the ammonium salts).

  [33] The change of colour is dependent in all probability on the
       combination with water, or according to others on polymeric
       transformation. It enables a solution of cobalt chloride to be
       used as sympathetic ink. If something be written with cobalt
       chloride on white paper, it will be invisible on account of the
       feeble colour of the solution, and when dry nothing can be
       distinguished. If, however, the paper be heated before the fire,
       the rose-coloured salt will be changed into a less hydrous blue
       salt, and the writing will become quite visible, but fade away
       when cool.

       The change of colour which takes place in solutions of CoCl_{2}
       under the influence not only of solution in water or alcohol, but
       also of a change of temperature, is a characteristic of all the
       halogen salts of cobalt. Crystalline iodide of cobalt,
       CoI_{2}6H_{2}O, gives a dark red solution between -22° and +20°;
       above +20° the solution turns brown and passes from olive to
       green, from +35° to 320° the solution remains green. According to
       Étard the change of colour is due to the fact that at first the
       solution contains the hydrate CoI_{2}H_{2}O, and that above 35° it
       contains CoI_{2}4H_{2}O. These hydrates can be crystallised from
       the solutions; the former at ordinary temperature and the latter
       on heating the solution. The intermediate olive colour of the
       solutions corresponds to the incipient decomposition of the
       hexahydrated salt and its passage into CoI_{2}4H_{2}O. A solution
       of the hexahydrated chloride of cobalt, CoCl_{2}6H_{2}O, is
       rose-coloured between -22° and +25°; but the colour changes
       starting from +25°, and passes through all the tints between red
       and blue right up to 50°; a true blue solution is only obtained at
       55° and remains up to 300°. This true blue solution contains
       another hydrate, CoCl_{2}2H_{2}O.

       The dependence between the solubility of the iodide and chloride
       of cobalt and the temperature is expressed by two almost straight
       lines corresponding to the hexa- and di-hydrates; the passage of
       the one into the other hydrate being expressed by a curve. The
       same character of phenomena is seen also in the variation of the
       vapour tension of solutions of chloride of cobalt with the
       temperature. We have repeatedly seen that aqueous solutions (for
       instance, Chapter XXII., Note 23 for Fe_{2}Cl_{6}) deposit
       different crystallo-hydrates at different temperatures, and that
       the amount of water in the hydrate decreases as the temperature
       _t_ rises, so that it is not surprising that CoCl_{2}2H_{2}O (or
       according to Potilitzin CoCl_{2}H_{2}O) should separate out above
       55° and CoCl_{2}6H_{2}O at 25° and below. Nor is it exceptional
       that the colour of a salt varies according as it contains
       different amounts of H_{2}O. But in this instance it is
       characteristic that the change of colour takes place in solution
       in the presence of an excess of water. This apparently shows that
       the actual solution may contain either CoCl_{2}6H_{2}O or
       CoCl_{2}2H_{2}O. And as we know that a solution may contain both
       metaphosphoric PHO_{3} and orthophosphoric acid H_{3}PO_{4} =
       HPO_{3} + H_{2}O, as well as certain other anhydrides, the
       question of the state of substances in solutions becomes still
       more complicated.

       Nickel sulphate crystallises from neutral solutions at a
       temperature of from 15° to 20° in _rhombic_ crystals containing
       7H_{2}O. Its form approaches very closely to that of the salts of
       zinc and magnesium. The planes of a vertical prism for magnesium
       salts are inclined at an angle of 90° 30´, for zinc salts at an
       angle of 91° 7´, and for nickel salts at an angle of 91° 10´. Such
       is also the form of the zinc and magnesium selenates and
       chromates. Cobalt sulphate containing 7 molecules of water is
       deposited in crystals of the _monoclinic_ system, like the
       corresponding salts of iron and manganese. The angle of a vertical
       prism for the iron salt = 82° 20´, for cobalt = 82° 22´, and the
       inclination of the horizontal pinacoid to the vertical prism for
       the iron salt = 99° 2´, and for the cobalt salt 99° 36´. All the
       isomorphous mixtures of the salts of magnesium, iron, cobalt,
       nickel and manganese have the same form if they contain 7 mol.
       H_{2}O and iron or cobalt predominate, whilst if there is a
       preponderance of magnesium, zinc, or nickel, the crystals have a
       rhombic form like magnesium sulphate. Hence these sulphates are
       _dimorphous_, but for some the one form is more stable and for
       others the other. Brooke, Moss, Mitscherlich, Rammelsberg, and
       Marignac have explained these relations. Brooke and Mitscherlich
       also supposed that NiSO_{4},7H_{2}O is not only capable of
       assuming these forms, but also that of the _tetragonal_ system,
       because it is deposited in this form from acid, and especially
       from slightly-heated solutions (30° to 40°). But Marignac
       demonstrated that the tetragonal crystals do not contain 7, but 6,
       molecules of water, NiSO_{4},6H_{2}O. He also observed that a
       solution evaporated at 50° to 70° deposits monoclinic crystals,
       but of a different form from ferrous sulphate,
       FeSO_{4},7H_{2}O--namely, the angle of the prism is 71° 52´, that
       of the pinacoid 95° 6´. This salt appears to be the same with 6
       molecules of water as the tetragonal. Marignac also obtained
       magnesium and zinc salts with 6 molecules of water by evaporating
       their solutions at a higher temperature, and these salts were
       found to be isomorphous with the monoclinic nickel salt. In
       addition to this it must be observed that the rhombic crystals of
       nickel sulphate with 7H_{2}O become turbid under the influence of
       heat and light, lose water, and change into the tetragonal salt.
       The monoclinic crystals in time also become turbid, and change
       their structure, so that the tetragonal form of this salt is the
       most stable. Let us also add that nickel sulphate in all its
       shapes forms very beautiful emerald green crystals, which, when
       heated to 230°, assume a dirty greenish-yellow hue and then
       contain one molecule of water.

       Klobb (1891) and Langlot and Lenoir obtained anhydrous CoSO_{4}
       and NiSO_{4} by igniting the hydrated salt with (NH_{4})_{2}SO_{4}
       until the ammonium salt had completely volatilised and decomposed.

       We may add that when equivalent aqueous solutions of NiX_{2}
       (green) and CoX_{2} (red) are mixed together they give an almost
       colourless (grey) solution, in which the green and red colour of
       the component parts disappears owing to the combination of the
       complementary colours.

       A double salt NiKF_{3} is obtained by heating NiCl_{2} with KFHF
       in a platinum crucible; KCoF_{3} is formed in a similar manner.
       The nickel salt occurs in fine green plates, easily soluble in
       water but scarcely soluble in ethyl and methyl alcohol. They
       decompose into green oxide of nickel and potassium fluoride when
       heated in a current of air. The analogous salt of cobalt
       crystallises in crimson flakes.

       If instead of potassium fluoride, CoCl_{2} or NiCl_{2} be fused
       with ammonium fluoride, they also form double salts with the
       latter. This gives the possibility of obtaining anhydrous
       fluorides NiF_{2} and CoF_{2}. Crystalline fluoride of nickel,
       obtained by heating the amorphous powder formed by decomposing the
       double ammonium salt in a stream of hydrofluoric acid, occurs in
       beautiful green prisms, sp. gr. 4·63, which are insoluble in
       water, alcohol, and ether; sulphuric, hydrochloric, and nitric
       acids also have no action upon them, even when heated; NiF_{2} is
       decomposed by steam, with the formation of black oxide, which
       retains the crystalline structure of the salt. Fluoride of cobalt,
       obtained as a rose-coloured powder by decomposing the double
       ammonium salt with the aid of heat in a stream of hydrofluoric
       acid, fuses into a ruby-coloured mass which bears distinct signs
       of a crystalline structure; sp. gr. 4·43. The molten salt only
       volatilises at about 1400°, which forms a clear distinction
       between CoF_{2} and the volatile NiF_{2}. Hydrochloric, sulphuric,
       and nitric acids act upon CoF_{2} even in the cold, although
       slowly, while when heated the reaction proceeds rapidly (Poulenc,
       1892).

If a solution of potassium hydroxide be added to a solution of a cobalt
salt, a blue precipitate of the basic salt will be formed. If a solution
of a cobalt salt be heated almost to the boiling-point, and the solution
be then mixed with a boiling solution of an alkali hydroxide, a _pink
precipitate of cobaltous hydroxide_, CoH_{2}O_{2}, will be formed. If air
be not completely excluded during the precipitation by boiling, the
precipitate will also contain brown cobaltic hydroxide formed by the
further oxidation of the cobaltous oxide.[34] Under similar circumstances
nickel salts form _a green precipitate of nickelous hydroxide_, the
formation of which is not hindered by the presence of ammonium salts, but
in that case only requires more alkali to completely separate the nickel.
The nickelous oxide obtained by heating the hydroxide, or from the
carbonate or nitrate, is a grey powder, easily soluble in acids and
easily reduced, but the same substance may be obtained in the crystalline
form as an ordinary product from the ores; it crystallises in regular
octahedra, with a metallic lustre, and is of a grey colour. In this state
the nickelous oxide almost resists the action of acids.[34 bis]

  [34] Hydrated suboxide of cobalt (de Schulten, 1889) is obtained in the
       following manner. A solution of 10 grams of CoCl_{2}6H_{2}O in 60
       c.c. of water is heated in a flask with 250 grams of caustic
       potash and a stream of coal gas is passed through the solution.
       When heated the hydrate of the suboxide of cobalt which separates
       out, dissolves in the caustic potash and forms a dark blue
       solution. This solution is allowed to stand for 24 hours in an
       atmosphere of coal gas (in order to prevent oxidation). The
       crystalline mass which separates out has a composition Co(OH)_{2},
       and to the naked eye appears as a violet powder, which is seen to
       be crystalline under the microscope. The specific gravity of this
       hydrate is 3·597 at 15°. It does not undergo change in the air;
       warm acetic acid dissolves it, but it is insoluble in warm and
       cold solutions of ammonia and sal-ammoniac.

  [34 bis] The following reaction may be added to those of the cobaltous
       and nickelous salts: potassium cyanide forms a precipitate with
       cobalt salts which is soluble in an excess of the reagent and
       forms a green solution. On heating this and adding a certain
       quantity of acid, a double _cobalt cyanide_ is formed which
       corresponds with potassium ferricyanide. Its formation is
       accompanied with the evolution of hydrogen, and is founded upon
       the property which cobalt has of oxidising in an alkaline
       solution, the development of which has been observed in such a
       considerable measure in the cobaltamine salts. The process which
       goes on here may be expressed by the following equation;
       CoC_{2}N_{2} + 4KCN first forms CoK_{4}C_{6}N_{6}, which salt with
       water, H_{2}O, forms potassium hydroxide, KHO, hydrogen, H, and
       the salt, K_{3}CoC_{6}N_{6}. Here naturally the presence of the
       acid is indispensable in consequence of its being required to
       combine with the alkali. From aqueous solutions this salt
       crystallises in transparent, hexagonal prisms of a yellow colour,
       easily soluble in water. The reactions of double decomposition,
       and even the formation of the corresponding acid, are here
       completely the same as in the case of the ferricyanide. If a
       nickelous salt be treated in precisely the same manner as that
       just described for a salt of cobalt, decomposition will occur.

It is interesting to note _the relation_ of the cobaltous and
nickelous hydroxides _to ammonia_; aqueous ammonia dissolves the
precipitate of cobaltous and nickelous hydroxide. The blue ammoniacal
solution of nickel resembles the same solution of cupric oxide, but has a
somewhat reddish tint. It is characterised by the fact that it dissolves
silk in the same way as the ammoniacal cupric oxide dissolves cellulose.
Ammonia likewise dissolves the precipitate of cobaltous hydroxide,
forming a brownish liquid, which becomes darker in air and finally
assumes a bright red hue, absorbing oxygen. The admixture of ammonium
chloride prevents the precipitation of cobalt salts by ammonia; when the
ammonia is added, a brown solution is obtained from which, as in the case
of the preceding solution, potassium hydroxide does not separate the
cobaltous oxide. Peculiar compounds are produced in this solution; they
are comparatively stable, containing ammonia and an excess of oxygen;
they bear the name cobaltoamine and cobaltiamine salts. They have been
principally investigated by Genth, Frémy, Jörgenson and others. Genth
found that when a cobalt salt, mixed with an excess of ammonium chloride,
is treated with ammonia and exposed to the air, after a certain lapse of
time, on adding hydrochloric acid and boiling, a red powder is
precipitated and the remaining solution contains an orange salt. The
study of these compounds led to the discovery of a whole series of
similar salts, some of which correspond with particular higher degrees of
oxidation of cobalt, which are described later.[35] Nickel does not
possess this property of absorbing the oxygen of the air when in an
ammoniacal solution. In order to understand this distinction, and in
general the relation of nickel, it is important to observe that cobalt
more easily forms a higher degree of oxidation--namely, _sesquioxide of
cobalt_, _cobaltic oxide_, Co_{2}O_{3}--than nickel, especially in the
presence of hypochlorous acid. If a solution of a cobalt salt be mixed
with barium carbonate and an excess of hypochlorous acid be added, or
chlorine gas be passed through it, then at the ordinary temperature on
shaking, the whole of the cobalt will be separated in the form of black
cobaltic oxide: 2CoSO_{4} + ClHO + 2BaCO_{3} = Co_{2}O_{3} + 2BaSO_{4} +
HCl + 2CO_{2}. Under these circumstances nickelous oxide does not
immediately form black sesquioxide, but after a considerable space of
time it also separates in the form of sesquioxide, Ni_{2}O_{3}, but
always later than cobalt. This is due to the relative difficulty of
further oxidation of the nickelous oxide. It is, however, possible to
oxidise it; if, for instance, the hydroxide NiH_{2}O_{2} be shaken in
water and chlorine gas be passed through it, then nickel chloride will be
formed, which is soluble in water, and insoluble nickelic oxide in the
form of a black precipitate: 3NiH_{2}O_{2} + Cl_{2} = NiCl_{2} +
Ni_{2}O_{3},3H_{2}O. Nickelic oxide may also be obtained by adding sodium
hypochlorite mixed with alkali to a solution of a nickel salt. Nickelic
and cobaltic hydrates are black. Nickelic oxide evolves oxygen with all
acids, and in consequence of this it is not separated as a precipitate in
the presence of acids; thus it evolves chlorine with hydrochloric acid,
exactly like manganese dioxide. When nickelic oxide is dissolved in
aqueous ammonia it liberates nitrogen, and an ammoniacal solution of
nickelous oxide is formed. When heated, nickelic oxide loses oxygen,
forming nickelous oxide. Cobaltic oxide, Co_{2}O_{3}, exhibits more
stability than nickelic oxide, and shows feeble basic properties; thus it
is dissolved in acetic acid without the evolution of oxygen.[35 bis] But
ordinary acids, especially on heating, evolve oxygen, forming a solution
of a cobaltous salt. The presence of a cobaltic salt in a solution of a
cobaltous salt may be detected by the brown colour of the solution and
the black precipitate formed by the addition of alkali, and also from the
fact that such solutions evolve chlorine when heated with hydrochloric
acid. Cobaltic oxide may not only be prepared by the above-mentioned
methods, but also by heating cobalt nitrate, after which a steel-coloured
mass remains which retains traces of nitric acid, but when heated further
to incandescence evolves oxygen, leaving a compound of cobaltic and
cobaltous oxides, similar to magnetic ironstone. Cobalt (but not nickel)
undoubtedly forms besides Co_{2}O_{3} a _dioxide_ CoO_{2}. This is
obtained[36] when the cobaltous oxide is oxidised by iodine or peroxide
of barium.[37]

  [35] The cobalt salts may be divided into at least the following
       classes, which repeat themselves for Cr, Ir, Rh (we shall not stop
       to consider the latter, particularly as they closely resemble the
       cobalt salts):--

       (_a_) _Ammonium cobalt salts_, which are simply direct compounds
       of the cobaltous salts CoX_{2} with ammonia, similar to various
       other compounds of the salts of silver, copper, and even calcium
       and magnesium, with ammonia. They are easily crystallised from an
       ammoniacal solution, and have a pink colour. Thus, for instance,
       when cobaltous chloride in solution is mixed with sufficient
       ammonia to redissolve the precipitate first formed, octahedral
       crystals are deposited which have a composition
       CoCl_{2},H_{2}O,6NH_{3}. These salts are nothing else but
       combinations with ammonia of crystallisation--if it may be so
       termed--likening them in this way to combinations with water of
       crystallisation. This similarity is evident both from their
       composition and from their capability of giving off ammonia at
       various temperatures. The most important point to observe is that
       all these salts contain 6 molecules of ammonia to 1 atom of
       cobalt, and this ammonia is held in fairly stable connection.
       Water decomposes these salts. (Nickel behaves similarly without
       forming other compounds corresponding to the true cobaltic.)

       (_b_) The solutions of the above-mentioned salts are rendered
       turbid by the action of the air; they absorb oxygen and become
       covered with a crust of _oxycobaltamine salts_. The latter are
       sparingly soluble in aqueous ammonia, have a brown colour, and are
       characterised by the fact that with warm water _they evolve
       oxygen_, forming salts of the following category: The nitrate may
       be taken as an example of this kind of salt; its composition is
       CoN_{2}O_{7},5NH_{3},H_{2}O. It differs from cobaltous nitrate,
       Co(NO_{3})_{2}, in containing an extra atom of oxygen--that is, it
       corresponds with cobalt dioxide, CoO_{2}, in the same way that the
       first salts correspond with cobaltous oxide; they contain 5, and
       not 6, molecules of ammonia, as if NH_{3} had been replaced by O,
       but we shall afterwards meet compounds containing either 5NH_{3}
       or 6NH_{3} to each atom of cobalt.

       (_c_) _The luteocobaltic salts_ are thus called because they have
       a yellow (luteus) colour. They are obtained from the salts of the
       first kind by submitting them in dilute solution to the action of
       the air; in this case salts of the second kind are not formed,
       because they are decomposed by an excess of water, with the
       evolution of oxygen and the formation of luteocobaltic salts. By
       the action of ammonia the salts of the fifth kind (roseocobaltic)
       are also converted into luteocobaltic salts. These last-named
       salts generally crystallise readily, and have a yellow colour;
       they are comparatively much more stable than the preceding ones,
       and even for a certain time resist the action of boiling water.
       Boiling aqueous potash liberates ammonia and precipitates hydrated
       cobaltic oxide, Co_{2}O_{3},3H_{2}O, from them. This shows that
       the luteocobaltic salts correspond with cobaltic oxide,
       Co_{2}O_{3}, and those of the second kind with the dioxide. When a
       solution of luteocobaltic sulphate,
       Co_{2}(SO_{4})_{3},12NH_{3},4H_{2}O, is treated with baryta,
       barium sulphate is precipitated, and the solution contains