Chapter XXII: Iron, Cobalt, and Nickel (2)

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.

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_