same acid. This barium salt may be prepared by the action of
barium chloride on the difficultly soluble silver permanganate,
AgMnO_{4}, which is precipitated on mixing a strong solution of
the potassium salt with silver nitrate. The solution of
permanganic acid forms a bright red liquid which reflects a dark
violet tint. A dilute solution has exactly the same colour as that
of the potassium salt. It deposits manganese dioxide when exposed
to the action of light, and also when heated above 60°, and this
proceeds the more rapidly the more dilute the solution. It shows
its oxidising properties in many cases, as already mentioned. Even
hydrogen gas is absorbed by a solution of permanganic acid; and
charcoal and sulphur are also oxidised by it, as they are by
potassium permanganate. This may be taken advantage of in
analysing gunpowder, because when it is treated with a solution of
potassium permanganate, all the sulphur is converted into
sulphuric acid and all the charcoal into carbonic anhydride.
Finely-divided platinum immediately decomposes permanganic acid.
With potassium iodide it liberates iodine (which may afterwards be
oxidised into iodic acid) (Mitscherlich, Fromherz, Aschoff, and
others). Ammonia does not form a corresponding salt with free
permanganic acid, because it is oxidised with evolution of
nitrogen. The oxidising action of permanganic acid in a strong
solution may be accompanied by flame and the formation of violet
fumes of permanganic acid; thus a strong solution of it takes fire
when brought into contact with paper, alcohol, alkaline sulphides,
fats, &c.
We may add that, according to Franke, 1 part of potassium
permanganate with 13 parts of sulphuric acid at 100° gives brown
crystals of the salt Mn_{2}(SO_{4})_{3},H_{2}SO_{4},4H_{2}O, which
gives a precipitate of hydrated manganese dioxide, H_{2}MnO_{3} =
MnO_{2}H_{2}O, when treated with water.
Spring, by precipitating potassium permanganate with sodium
sulphite and washing the precipitate by decantation, obtained a
soluble colloidal manganese oxide, whose composition was the mean
between Mn_{2}O_{3} and MnO_{2}--namely,
Mn_{2}O_{3},4(MnO_{2}H_{2}O).
The oxidising action of KMnO_{4}, like all other chemical reactions, is
not accomplished instantaneously, but only gradually. And, as the course
of the reaction is here easily followed by determining the amount of salt
unchanged in a sample taken at a given moment,[25] the oxidising reaction
of potassium permanganate, in an acid liquid, was employed by Harcourt
and Esson (1865) as one of the first cases for the investigation of the
laws of the _rate of chemical change_[26] as a subject of great
importance in chemical mechanics. In their experiments they took oxalic
acid, C_{2}H_{2}O_{4}, which in oxidising gives carbonic anhydride,
whilst, with an excess of sulphuric acid, the potassium permanganate is
converted into manganous sulphate, MnSO_{4}, so that the ultimate
oxidation will be expressed by the equation: 5C_{2}H_{2}O_{4} + 2MnKO_{4}
+ 3H_{2}SO_{4} = 10CO_{2} + K_{2}SO_{4} + 2MnSO_{4} + 8H_{2}O. The
influence of the relative amount of sulphuric acid is seen from the
annexed table, which gives the measure of reaction _p_ per 100 parts of
potassium permanganate, taken four minutes after mixing, using n
molecules of sulphuric acid, H_{2}SO_{4}, per 2KMnO_{4} +
5C_{2}H_{2}O_{4}:
_n_ = 2 4 6 8 12 16 22
_p_ = 22 36 51 63 77 86 92
showing that in a given time (4 minutes) the oxidation is the more
perfect the greater the amount of sulphuric acid taken for given amounts
of KMnO_{4} and C_{2}H_{2}O_{4}. It is obvious also that the temperature
and relative amount of every one of the acting and resulting substances
should show its influence on the relative velocity of reaction; thus, for
instance, direct experiment showed the influence of the admixture of
manganous sulphate. When a large proportion of oxalic acid (108
molecules) was taken to a large mass of water and to 2 molecules of
permanganate 14 molecules of manganous sulphate were added, the quantity
x of the potassium permanganate acted on (in percentages of the potassium
permanganate taken) in t minutes (at 16°) was as follows:
_t_ = 2 5 8 11 14 44 47 53 61 68
_x_ = 5·2 12·1 18·7 25·1 31·3 68·4 71·7 75·8 79·8 83·0
These figures show that the rate of reaction--that is, the quantity of
permanganate changed in one minute--decreases proportionally to the
decrease in the amount of unchanged potassium permanganate. At the
commencement, about 2·6 per cent. of the salt taken was decomposed in the
course of one minute, whilst after an hour the rate was about 0·5 per
cent. The same phenomena are observed in every case which has been
investigated, and this branch of theoretical or physical chemistry, now
studied by many,[27] promises to explain the course of chemical
transformations from a fresh point of view, which is closely allied to
the doctrine of affinity, because the rate of reaction, without doubt, is
connected with the magnitude of the affinities acting between the
reacting substances.
[25] For rapid and accurate determinations of this kind, advantage is
taken of those methods of chemical analysis which are known as
'titrations' (volumetric analysis), and consist in measuring the
volume of solutions of known strength required for the complete
conversion of a given substance. Details respecting the theory and
practice of titration, in which potassium permanganate is very
frequently employed, must be looked for in works on analytical
chemistry.
[26] The measurements of velocity and acceleration serve for
determining the measure of forces in mechanics, but in that case
the velocities are magnitudes of length or paths passed over in a
unit of time. The velocity of chemical change embodies a
conception of quite another kind. In the first place, the
velocities of reactions are magnitudes of the masses which have
entered into chemical transformations; in the second place, these
velocities can only be relative quantities. Hence the conception
of 'velocity' has quite a different meaning in chemistry from what
it has in mechanics. Their only common factor is time. If _dt_ be
the increment of time and _dx_ the quantity of a substance changed
in this space of time, then the fraction (or quotient) _dx/dt_
will express the rate of the reaction. The natural conclusion,
come to both by Harcourt and Esson, and previously to them (1850)
by Wilhelmj (who investigated the rate of conversion, or
inversion, of sugar in its passage into glucose), consists in
establishing that this velocity is proportional to the quantity of
substances still unchanged--_i.e._ that _dx/dt_ = C(A - _x_),
where C is a constant coefficient of proportionality, and where A
is the quantity of a substance taken for reaction at the moment
when _t_ = 0 and _x_ = 0--that is, at the beginning of the
experiment, from which the time _t_ and quantity _x_ of substance
changed is counted. On integrating the preceding equation we
obtain log(A/(A - _x_)) = _kt_, where _k_ is a new constant, if we
take ordinary (and not natural) logarithms. Hence, knowing A, _x_,
and _t_, for each reaction, we find _k_, and it proves to be a
constant quantity. Thus from the figures cited in the text for the
reaction 2KMnO_{4} + 108C_{2}H_{2}O_{4} + 14MnSO_{4}, it may be
calculated that _k_ = 0·0114; for example, _t_ = 44, _x_ = 68·4 (A
= 100), whence _kt_ = 0·5004 and _k_ = 0·0114, (_see also_ Chapter
XIV., Note 3, and Chapter XVII., Note 25 bis).
[27] The researches made by Hood, Van't Hoff, Ostwald, Warder,
Menschutkin, Konovaloff, and others have a particular significance
in this direction. Owing to the comparative novelty of this
subject, and the absence of applicable as well as indubitable
deductions, I consider it impossible to enter into this province
of theoretical chemistry, although I am quite confident that its
development should lead to very important results, especially in
respect to chemical equilibria, for Van't Hoff has already shown
that the limit of reaction in reversible reactions is determined
by the attainment of equal velocities for the opposite reactions.
CHAPTER XXII
IRON, COBALT, AND NICKEL
Judging from the atomic weights, and the forms of the higher oxides of
the elements already considered, it is easy to form an idea of the seven
groups of the periodic system. Such are, for instance, the typical series
Li, Be, B, C, N, O, F, or the third series, Na, Mg, Al, Si, P, S, Cl. The
seven usual types of oxides from R_{2}O to R_{2}O_{7} correspond with
them (Chapter XV.) The position of the eighth group is quite separate,
and is determined by the fact that, as we have already seen, in each
group of metals having a greater atomic weight than potassium a
distinction ought to be made between the elements of the even and uneven
series. The series of even elements, commencing with a strikingly
alkaline element (potassium, rubidium, cæsium), together with the uneven
series following it, and concluding with a haloid (chlorine, bromine,
iodine), forms a large period, the properties of whose members repeat
themselves in other similar periods. The elements of the eighth group are
situated between the elements of the even series and the elements of the
uneven series following them. And for this reason elements of the eighth
group are found in the middle of each large period. The properties of the
elements belonging to it, in many respects independent and striking, are
shown with typical clearness in the case of iron, the well-known
representative of this group.
_Iron_ is one of those elements which are not only widely diffused in the
crust of the earth, but also throughout the entire universe. Its oxides
and their various compounds are found in the most diverse portions of the
earth's crust; but here iron is always found combined with some other
element. Iron is not found on the earth's surface in a free state,
because it easily oxidises under the action of air. It is occasionally
found in the native state in meteorites, or aerolites, which fall upon
the earth.
_Meteoric iron_ is formed outside the earth.[1] Meteorites are
fragments which are carried round the sun in orbits, and fall upon the
earth when coming into proximity with it during their motion in space.
The meteoric dust, on passing through the upper parts of the atmosphere,
and becoming incandescent from friction with the gases, produces that
phenomenon which is familiar under the name of falling stars.[2] Such is
the doctrine concerning meteorites, and therefore the fact of their
containing rocky (siliceous) matter and metallic iron shows that outside
the earth the elements and their aggregation are in some degree the same
as upon the earth itself.
[1] The composition of meteoric iron is variable. It generally contains
nickel, phosphorus, carbon, &c. The schreibersite of meteoric
stones contains Fe_{4}Ni_{2}P.
[2] Comets and the rings of Saturn ought now to be considered as
consisting of an accumulation of such meteoric cosmic particles.
Perhaps the part played by these minute bodies scattered throughout
space is much more important in the formation of the largest
celestial bodies than has hitherto been imagined. The investigation
of this branch of astronomy, due to Schiaparelli, has a bearing on
the whole of natural science.
The question arises as to why the iron in meteorites is in a free
state, whilst on earth it is in a state of combination. Does not
this tend to show that the condition of our globe is very different
from that of the rest? My answer to this question has been already
given in Volume I. p. 377, Note 57. It is my opinion that inside
the earth there is a mass similar in composition to
meteorites--that is, containing rocky matter and metallic iron,
partly carburetted. In conclusion, I consider it will not be out of
place to add the following explanations. According to the theory of
the distribution of pressures (see my treatise, _On Barometrical
Levelling_, 1876, pages 48 _et seq._) in an atmosphere of mixed
gases, it follows that two gases, whose densities are _d_ and
_d__{1}, and whose relative quantities or partial pressures at a
certain distance from the centre of gravity are _h_ and _h__{1},
will, when at a greater distance from the centre of attraction,
present a different ratio of their masses _x_ : _x__{1}--that is,
of their partial pressures--which may be found by the equation
_d__{1}(log(_h_) - log(_x_)) = _d_(log(_h__{1}) - log(_x__{1})).
If, for instance, _d_ : _d__{1} = 2 : 1, and _h_ = _h__{1} (that is
to say, the masses are equal at the lower height) = 1000, then when
_x_ = 10 the magnitude of _x__{1} will not be 10 (_i.e._ the mass
of a gas at a higher level whose density = 1 will not be equal to
the mass of a gas whose density = 2, as was the case at a lower
level), but much greater--namely, _x__{1} = 100--that is, the
lighter gas will predominate over a heavier one at a higher level.
Therefore, when the whole mass of the earth was in a state of
vapour, the substances having a greater vapour density accumulated
about the centre and those with a lesser vapour density at the
surface. And as the vapour densities depend on the atomic and
molecular weights, those substances which have small atomic and
molecular weights ought to have accumulated at the surface, and
those with high atomic and molecular weights, which are the least
volatile and the easiest to condense, at the centre. Thus it
becomes apparent why such light elements as hydrogen, carbon,
nitrogen, oxygen, sodium, magnesium, aluminium, silicon,
phosphorus, sulphur, chlorine, potassium, calcium, and their
compounds predominate at the surface and largely form the earth's
crust. There is also now much iron in the sun, as spectrum analysis
shows, and therefore it must have entered into the composition of
the earth and other planets, but would have accumulated at the
centre, because the density of its vapour is certainly large and it
easily condenses. There was also oxygen near the centre of the
earth, but not sufficient to combine with the iron. The former, as
a much lighter element, principally accumulated at the surface,
where we at the present time find all oxidised compounds and even a
remnant of free oxygen. This gives the possibility not only of
explaining in accordance with cosmogonic theories the predominance
of oxygen compounds on the surface of the earth, with the
occurrence of unoxidised iron in the interior of the earth and in
meteorites, but also of understanding why the density of the whole
earth (over 5) is far greater than that of the rocks (1 to 3)
composing its crust. And if all the preceding arguments and
theories (for instance the supposition that the sun, earth, and all
the planets were formed of an elementary homogeneous mass, formerly
composed of vapours and gases) be true, it must be admitted that
the interior of the earth and other planets contains metallic
(unoxidised) iron, which, however, is only found on the surface as
aerolites. And then assuming that aerolites are the fragments of
planets which have crumbled to pieces so to say during cooling
(this has been held to be the case by astronomers, judging from the
paths of aerolites), it is readily understood why they should be
composed of metallic iron, and this would explain its occurrence in
the depths of the earth, which we assumed as the basis of our
theory of the formation of naphtha (Chapter VIII., Notes 57-60).
The most widely diffused terrestrial compound of iron is iron bisulphide,
FeS_{2}, or _iron pyrites_. It occurs in formations of both aqueous and
igneous origin, and sometimes in enormous masses. It is a substance
having a greyish-yellow colour, with a metallic lustre, and a specific
gravity of 5·0; it crystallises in the regular system.[2 bis]
[2 bis] Immense deposits of iron pyrites are known in various parts of
Russia. On the river Msta, near Borovitsi, thousands of tons are
yearly collected from the detritus of the neighbouring rocks. In
the Governments of Toula, Riazan, and in the Donets district
continuous layers of pyrites occur among the coal seams. Very thick
beds of pyrites are also known in many parts of the Caucasus. But
the deposits of the Urals are particularly vast, and have been
worked for a long time. Amongst these I will only indicate the
deposits on the Soymensky estate near the Kishteimsky works; the
Kaletinsky deposits near the Virhny-Isetsky works (containing 1-2
p.c. Cu); on the banks of the river Koushaivi near Koushvi (3-5
p.c. Cu), and the deposits near the Bogoslovsky works (3-5 p.c.
Cu). Iron pyrites (especially that containing copper which is
extracted after roasting) is now chiefly employed for roasting, as
a source of SO_{2}, for the manufacture of chamber sulphuric acid
(Vol. I. p. 291), but the remaining oxide of iron is perfectly
suitable for smelting into pig iron, although it gives a sulphurous
pig iron (the sulphur may be easily removed by subsequent
treatment, especially with the aid of ferro-manganese in Bessemer's
process). The great technical importance of iron pyrites leads to
its sometimes being imported from great distances; for instance,
into England from Spain. Besides which, when heated in closed
retorts FeS_{2} gives sulphur, and if allowed to oxidise in damp
air, green vitriol, FeSO_{4}.
The oxides are the principal ores used for producing metallic iron. The
majority of the ores contain ferric oxide, Fe_{2}O_{3}, either in a free
state or combined with water, or else in combination with ferrous oxide,
FeO. The species and varieties of iron ores are numerous and diverse.
Ferric oxide in a separate form appears sometimes as crystals of the
rhombohedric system, having a metallic lustre and greyish steel colour;
they are brittle, and form a red powder, specific gravity about 5·25.
Ferric oxide in type of oxidation and properties resembles alumina; it
is, however, although with difficulty, soluble in acids even when
anhydrous. The crystalline oxide bears the name of _specular iron ore_,
but ferric oxide most often occurs in a non-crystalline form, in masses
having a red fracture, and is then known as _red hæmatite_. In this form,
however, it is rather a rare ore, and is principally found in veins. The
hydrates of ferric oxide, ferric hydroxides,[3] are most often found in
aqueous or stratified formations, and are known as _brown hæmatites_;
they generally have a brown colour, form a yellowish-brown powder, and
have no metallic lustre but an earthy appearance. They easily dissolve in
acids and diffuse through other formations, especially clays (for
instance, ochre); they sometimes occur in reniform and similar masses,
evidently of aqueous origin. Such are, for instance, the so-called bog or
lake and peat ores found at the bottom of marshes and lakes, and also
under and in peat beds. This ore is formed from water containing ferrous
carbonate in solution, which, after absorbing oxygen, deposits ferric
hydroxide. In rivers and springs, iron is found in solution as ferrous
carbonate through the agency of carbonic acid: hence the existence of
chalybeate springs containing FeCO_{3}. This ferrous carbonate, or
_siderite_, is either found as a non-crystalline product of evidently
aqueous origin, or as a crystalline spar called _spathic iron ore_. The
reniform deposits of the former are most remarkable; they are called
spherosiderites, and sometimes form whole strata in the jurassic and
carboniferous formations. _Magnetic iron ore_, Fe_{3}O_{4} =
FeO,Fe_{2}O_{3}, in virtue of its purity and practical uses, is a very
important ore; it is a compound of the ferrous and ferric oxides, is
naturally magnetic, has a specific gravity of 5·1, crystallises in
well-formed crystals of the regular system, is with difficulty soluble in
acids, and sometimes forms enormous masses, as, for instance, Mount
Blagodat in the Ural. However, in most cases--for instance, at
Korsak-Mogila (to the north of Berdiansk and Nogaiska, near the Sea of
Azov), or at Krivoi Rog (to the west of Ekaterinoslav)--the magnetic iron
ore is mixed with other iron ores. In the Urals, the Caucasus (without
mentioning Siberia), and in the districts adjoining the basin of the Don,
Russia possesses the richest iron ores in the world. To the south of
Moscow, in the Governments of Toula and Nijninovgorod, in the Olonetz
district, and in the Government of Orloffsky (near Zinovieff in the
district of Kromsky), and in many other places, there are likewise
abundant supplies of iron ores amongst the deposited aqueous formations;
the siderite of Orloffsky, for instance, is distinguished by its great
purity.[4]
[3] The hydrated ferric oxide is found in nature in a dual form. It is
somewhat rarely met with in the form of a crystalline mineral
called _göthite_, whose specific gravity is 4·4 and composition
Fe_{2}H_{3}O_{4}, or FeHO_{2}--that is, one of oxide of iron to one
of water, Fe_{2}O_{3},H_{2}O; frequently found as brown ironstone,
forming a dense mass of fibrous, reniform deposits containing
2Fe_{2}O_{3},3H_{2}O--that is, having a composition
Fe_{4}H_{6}O_{9}. In bog ore and other similar ores we most often
find a mixture of this hydrated ferric oxide with clay and other
impurities. The specific gravity of such formations is rarely as
high as 4·0.
[4] The ores of iron, similarly to all substances extracted from veins
and deposits, are worked according to mining practice by means of
vertical, horizontal, or inclined shafts which reach and penetrate
the veins and strata containing the ore deposits. The mass of ore
excavated is raised to the surface, then sorted either by hand or
else in special sorting apparatus (generally acting with water to
wash the ore), and is subjected to roasting and other treatment. In
every case the ore contains foreign matter. In the extraction of
iron, which is one of the cheapest metals, the dressing of an ore
is in most cases unprofitable, and only ores rich in metal are
worked--namely, those containing at least 20 p.c. It is often
profitable to transport very rich and pure ores (with as much as 70
p.c. of iron) from long distances. The details concerning the
working and extraction of metals will be found in special treatises
on metallurgy and mining.
Iron is also found in the form of various other compounds--for instance,
in certain silicates, and also in some phosphates; but these forms are
comparatively rare in nature in a pure state, and have not the industrial
importance of those natural compounds of iron previously mentioned. In
small quantities iron enters into the composition of every kind of _soil_
and all rocky formations. As ferrous oxide, FeO, is isomorphous with
magnesia, and ferric oxide, Fe_{2}O_{3}, with alumina, isomorphous
substitution is possible here, and hence minerals are not unfrequently
found in which the quantity of iron varies considerably; such, for
instance, are pyroxene, amphibole, certain varieties of mica, &c.
Although much iron oxide is deleterious to the growth of vegetation,
still plants do not flourish without iron; it enters as an indispensable
component into the composition of all higher _organisms_; in the ash of
plants we always find more or less of its compounds. It also occurs in
blood, and forms one of the colouring matters in it; 100 parts of the
blood of the highest organisms contain about 0·05 of iron.
The _reduction_ of the ores of iron into metallic iron is in principle
very simple, because when the oxides of iron are strongly heated with
charcoal, hydrogen, carbonic oxide, and other reducing agents,[5] they
easily give metallic iron. But the matter is rendered more difficult by
the fact that the iron does not melt at the heat developed by the
combustion of the charcoal, and therefore it does not separate from those
mechanically mixed impurities which are found in the iron ore. This is
obviated by the following very remarkable property of iron: at a high
temperature it is capable of combining with a small quantity (from 2 to 5
p.c.) of carbon, and then forms _cast iron_, which easily _melts_ in the
heat developed by the combustion of charcoal in air. For this reason
metallic iron is not obtained directly from the ore, but is only formed
after the further treatment of the cast iron; the first product extracted
from the ore being cast iron. The fused mass disposes itself in the
furnace below the slag--that is, the impurities of the ore fused by the
heat of the furnace. If these impurities did not fuse they would block up
the furnace in which the ore was being smelted, and the continuous
smelting of the cast iron would not be possible;[6] it would be necessary
periodically to cool the furnace and heat it up again, which means a
wasteful expenditure of fuel, and hence in the production of cast iron,
the object in view is to obtain all the earthy impurities of the ore in
the shape of a fused mass or slag. Only in rare cases does the ore itself
form a mass which fuses at the temperature employed, and these cases are
objectionable if much iron oxide is carried away in the slag. The
impurities of the ores most often consist of certain mixtures--for
instance, a mixture of clay and sand, or a mixture of limestone and clay,
or quartz, &c. These impurities do not separate of themselves, or do not
fuse. The difficulty of the industry lies in forming an easily-fusible
slag, into which the whole of the foreign matter of the ore would pass
and flow down to the bottom of the furnace above the heavier cast iron.
This is effected by mixing certain _fluxes_ with the ore and charcoal. A
flux is a substance which, when mixed with the foreign matter of the ore,
forms a fusible vitreous mass or slag. The flux used for silica is
limestone with clay; for limestone a definite quantity of silica is used,
the best procedure having been arrived at by experiment and by long
practice in iron smelting and other metallurgical processes.[7]
[5] The reduction of iron oxides by hydrogen belongs to the order of
reversible reactions (Chapter II.), and is therefore determined by
a limit which is here expressed by the attainment of the same
pressure as in the case where hydrogen acts on iron oxides, and as
in the case where (at the same temperature) water is decomposed by
metallic iron. The calculations referring to this matter were made
by Henri Sainte-Claire Deville (1870). Spongy iron was placed in a
tube having a temperature _t_, one end of which was connected with
a vessel containing water at 0° (vapour tension = 4·6 mm.) and the
other end with a mercury pump and pressure gauge which determined
the limiting tension attained by the dry hydrogen _p_ (subtracting
the tension of the water vapour from the tension observed). A tube
was then taken containing an excess of iron oxide. It was filled
with hydrogen, and the tension _p__{1} observed of the residual
hydrogen when the water was condensed at 0°.
_t_ = 200° 440° 860° 1040°
_p_ = 95·9 25·8 12·8 9·2 mm.
_p_{1} = -- -- 12·8 9·4 mm.
The equality of the pressure (tension) of the hydrogen in the two
cases is evident. The hydrogen here behaves like the vapour of iron
or of its oxide.
By taking ferric oxide, Fe_{2}O_{3}, Moissan observed that at 350°
it passed into magnetic oxide, Fe_{3}O_{4}, at 500° into ferrous
oxide, FeO, and at 600° into metallic iron. Wright and Luff (1878),
whilst investigating the reduction of oxides, found that (_a_) the
temperature of reaction depends on the condition of the oxide
taken--for instance, precipitated ferric oxide is reduced by
hydrogen at 85°, that obtained by oxidising the metal or from its
nitrate at 175°; (_b_) when other conditions are the same, the
reduction by carbonic oxide commences earlier than that by
hydrogen, and the reduction by hydrogen still earlier than that by
charcoal; (_c_) the reduction is effected with greater facility
when a greater quantity of heat is evolved during the reaction.
Ferric oxide obtained by heating ferrous sulphate to a red heat
begins to be reduced by carbonic oxide at 202°, by hydrogen at
260°, by charcoal at 430°, whilst for magnetic oxide, Fe_{3}O_{4},
the temperatures are 200°, 290°, and 450° respectively.
[6] The primitive methods of iron manufacture were conducted by
intermittent processes in hearths resembling smiths' fires. As
evidenced by the uninterrupted action of the steam boiler, or the
process of lime burning, and the continuous preparation and
condensation of sulphuric acid or the uninterrupted smelting of
iron, every industrial process becomes increasingly profitable and
complete under the condition of the continuous action, as far as
possible, of all agencies concerned in the production. This
continuous method of production is the first condition for the
profitable production on the large scale of nearly all industrial
products. This method lessens the cost of labour, simplifies the
supervision of the work, renders the product uniform, and
frequently introduces a very great economy in the expenditure of
fuel and at the same time presents the simplicity and perfection of
an equilibrated system. Hence every manufacturing operation should
be a continuous one, and the manufacture of pig iron and sulphuric
acid, which have long since become so, may be taken as examples in
many respects. A study of these two manufactures should form the
commencement of an acquaintance with all the contemporary methods
of manufacturing both from a technical and economical point of
view.
[7] The composition of slag suitable for iron smelting most often
approaches the following: 50 to 60 p.c. SiO_{2}, 5 to 20
Al_{2}O_{3}, the rest of the mass consisting of MgO, CaO, MnO, FeO.
Thus the most fusible slag (according to the observations of
Bodeman) contains the alloy Al_{2}O_{3},4CaO,7SiO_{2}. On altering
the quantity of magnesia and lime, and especially of the alkalis
(which increases the fusibility) and of silica (which decreases
it), the temperature of fusion changes with the relation between
the total quantity of oxygen and that in the silica. Slags of the
composition RO,SiO_{2} are easily fusible, have a vitreous
appearance, and are very common. Basic slags approach the
composition 2RO,SiO_{2}. Hence, knowing the composition and
quantity of the foreign matter in the ore, it is at once easy to
find the quantity and quality of the flux which must be added to
form a suitable slag. The smelting of iron is rendered more complex
by the fact that the silica, SiO_{2}, which enters into the slag
and fluxes is capable of forming a slag with the iron oxides. In
order that the least quantity of iron may pass into the slag, it is
necessary for it to be reduced before the temperature is attained
at which the slags are formed (about 1000°), which is effected by
reducing the iron, not with charcoal itself, but with carbonic
oxide. From this it will be understood how the progress of the
whole treatment may be judged by the properties of the slags.
Details of this complicated and well-studied subject will be found
in works on metallurgy.
Thus the following materials have to be introduced into the furnace
where the smelting of the iron ore is carried on: (1) the iron ore,
composed of oxide of iron and foreign matter; (2) the flux required to
form a fusible slag with the foreign matter; (3) the carbon which is
necessary (_a_) for reducing, (_b_) for combining with the reduced iron
to form cast iron, (_c_) principally for the purpose of combustion and
the heat generated thereby, necessary not only for reducing the iron and
transforming it into cast iron, but also for melting the slag, as well as
the cast iron--and (4) the air necessary for the combustion of the
charcoal. The air is introduced after a preparatory heating in order to
economise fuel and to obtain the highest temperature. The air is forced
in under pressure by means of a special blast arrangement. This permits
of an exact regulation of the heat and rate of smelting. All these
component parts necessary for the smelting of iron must be contained in a
vertical, that is, _shaft furnace_, which at the base must have a
receptacle for the accumulation of the slag and cast iron formed, in
order that the operation may proceed without interruption. The walls of
such a furnace ought to be built of fireproof materials if it be designed
to serve for the continuous production of cast iron by charging the ore,
fuel, and flux into the mouth of the furnace, forcing a blast of air into
the lower part, and running out the molten iron and slag from below. The
whole operation is conducted in furnaces known as _blast furnaces_. The
annexed illustration, fig. 93 (which is taken by kind permission from
Thorpe's Dictionary of Applied Chemistry), represents the vertical
section of such a furnace. These furnaces are generally of large
dimensions--varying from 50 to 90 feet in height. They are sometimes
built against rising ground in order to afford easy access to the top
where the ore, flux, and charcoal or coke are charged.[8]
[8] The section of a blast furnace is represented by two truncated
cones joined at their bases, the upper cone being longer than the
lower one; the lower cone is terminated by the hearth, or almost
cylindrical cavity in which the cast iron and slag collect, one
side being provided with apertures for drawing off the iron and
slag. The air is blown into the blast furnace through special
pipes, situated over the hearth, as shown in the section. The air
previously passes through a series of cast-iron pipes, heated by
the combustion of the carbonic oxide obtained from the upper parts
of the furnace, where it is formed as in a 'gas-producer.' The
blast furnace acts continuously until it is worn out; the iron is
tapped off twice a day, and the furnace is allowed to cool a little
from time to time so as not to be spoilt by the increasing heat,
and to enable it to withstand long usage.
Blast furnaces worked with charcoal fuel are not so high, and in
general give a smaller yield than those using coke, because the
latter are worked with heavier charges than those in which charcoal
is employed. Coke furnaces yield 20,000 tons and over of pig iron a
year. In the United States there are blast furnaces 30 metres high,
and upwards of 600 cubic metres capacity, yielding as much as
130,000 tons of pig iron, requiring a blast of about 750 cubic
metres of air per minute, heated to 600°, and consuming about 0·85
part of coke per 1 part of pig iron produced. At the present time
the world produces as much as 30 million tons of pig iron a year,
about 9/10 of which is converted into wrought iron and steel. The
chief producers are the United States (about 10 million tons a
year) and England (about 9 million tons a year); Russia yields
about 1-1/5 million tons a year. The world's production has doubled
during the last 20 years, and in this respect the United States
have outrun all other countries. The reason of this increase of
production must be looked for in the increased demand for iron and
steel for railway purposes, for structures (especially
ship-building), and in the fact that: (_a_) the cost of pig iron
has fallen, thanks to the erection of large furnaces and a fuller
study of the processes taking place in them, and (_b_) that every
kind of iron ore (even sulphurous and phosphoritic) can now be
converted into a homogeneous steel.
[Illustration: FIG. 93.--Vertical section of a modern Cleveland
blast furnace capable of producing 300 to 1,000 tons of pig iron
weekly. The outer casing is of riveted iron plates, the furnace
being lined with refractory fire-brick. It is closed at the top by
a 'cap and cone' arrangement, by means of which the charge can be
fed into the furnace at suitable intervals by lowering the moveable
cone.]
In order to more thoroughly grasp the chemical process which takes
place in blast furnaces, it is necessary to follow the course of
the material charged in at the top and of the air passing through
the furnace. From 50 to 200 parts of carbon are expended on 100
parts of iron. The ore, flux, and coke are charged into the top of
the furnace, in layers, as the cast iron is formed in the lower
parts and flowing down to the bottom causes the whole contents of
the furnace to subside, thus forming an empty space at the top,
which is again filled up with the afore-mentioned mixture. During
its downward course this mixture is subjected to increasing heat.
This rise of temperature first drives off the moisture of the ore
mixture, and then leads to the formation of the products of the dry
distillation of coal or charcoal. Little by little the subsiding
mass attains a temperature at which the heated carbon reacts with
the carbonic anhydride passing upwards through the furnace and
transforms it into carbonic oxide. This is the reason why carbonic
anhydride is not evolved from the furnace, but only carbonic oxide.
As regards the ore itself, on being heated to about 600° to 800° it
is reduced at the expense of the carbonic oxide ascending the
furnace, and formed by the contact of the carbonic anhydride with
the incandescent charcoal, so that the reduction in the blast
furnace is without doubt brought about _by_ the formation and
decomposition of _carbonic oxide_ and not by carbon itself--thus,
Fe_{2}O_{3} + 3CO = Fe_{2} + 3CO_{2}. The reduced iron, on further
subsidence and contact with carbon, forms cast iron, which flows to
the bottom of the furnace. In these lower layers, where the
temperature is highest (about 1,300°), the foreign matter of the
ore finally forms slag, which also is fusible, with the aid of
fluxes. The air blown in from below, through the so-called
_tuyeres_, encounters carbon in the lower layers of the furnace,
and burns it, converting it into carbonic anhydride. It is evident
that this develops the highest temperature in these lower layers of
the furnace, because here the combustion of the carbon is effected
by heated and compressed air. This is very essential, for it is by
virtue of this high temperature that the process of forming the
slag and of forming and fusing the cast iron are effected
simultaneously in these lower portions of the furnace. The carbonic
acid formed in these parts rises higher, encounters incandescent
carbon, and forms with it carbonic oxide. This heated carbonic
oxide acts as a reducing agent on the iron ore, and is reconverted
by it into carbonic anhydride; this gas meets with more carbon, and
again forms carbonic oxide, which again acts as a reducing agent.
The final transformation of the carbonic anhydride into carbonic
oxide is effected in those parts of the furnace where the reduction
of the oxides of iron does not take place, but where the
temperature is still high enough to reduce the carbonic anhydride.
The ascending mixture of carbonic oxide and nitrogen, CO_{2}, &c.,
is then withdrawn through special lateral apertures formed in the
upper cold parts of the furnace walls, and is conducted through
pipes to those stoves which are used for heating the air, and also
sometimes into other furnaces used for the further processes of
iron manufacture. The fuel of blast furnaces consists of wood
charcoal (this is the most expensive material, but the pig iron
produced is the purest, because charcoal does not contain any
sulphur, while coke does), anthracite (for instance, in
Pennsylvania, and in Russia at Pastouhoff's works in the Don
district), coke, coal, and even wood and peat. It must be borne in
mind that the utilisation of naphtha and naphtha refuse would
probably give very profitable results in metallurgical processes.
The process just described is accompanied by a series of other
processes. Thus, for instance, in the blast furnace a considerable
quantity of cyanogen compounds are formed. This takes place because
the nitrogen of the air blast comes into contact with incandescent
carbon and various alkaline matters contained in the foreign matter
of the ores. A considerable quantity of potassium cyanide is formed
when wood charcoal is employed for iron smelting, as its ash is
rich in potash.
The _cast iron_ formed in blast furnaces is not always of the same
quality. When slowly cooled it is soft, has a grey colour, and is not
completely soluble in acids. When treated with acids a residue of
graphite remains; it is known as _grey_ or soft cast iron. This is the
general form of the ordinary cast iron used for casting various objects,
because in this state it is not so brittle as in the shape of _white cast
iron_, which does not leave particles of graphite when dissolved, but
yields its carbon in the form of hydrocarbons. This white cast iron is
characterised by its whitish-grey colour, dull lustre, the crystalline
structure of its fracture (more homogeneous than that of grey iron), and
such hardness that a file will hardly cut it. When white cast iron is
produced (from manganese ore) at high temperatures (and with an excess of
lime), and containing little sulphur and silica but a considerable amount
of carbon (as much as 5 p.c.), it acquires a coarse crystalline structure
which increases in proportion to the amount of manganese, and it is then
known under the name of 'spiegeleisen' (and 'ferro-manganese').[9]
[9] The specific gravity of white cast iron is about 7·5. Grey cast
iron has a much lower specific gravity, namely, 7·0. Grey cast iron
generally contains less manganese and more silica than white; but
both contain from 2 to 3 p.c. of carbon. The difference between the
varieties of cast iron depends on the condition of the carbon which
enters into the composition of the iron. In white cast iron the
carbon is in combination with the iron--in all probability, as the
compound CFe_{4} (Abel and Osmond and others extracted this
compound, which is sometimes called 'carbide,' from tempered steel,
which stands to unannealed steel as white cast iron does to grey),
but perhaps in the state of an indefinite chemical compound
resembling a solution. In any case the compound of the iron and
carbon in white cast iron is chemically very unstable, because when
slowly cooled it decomposes, with separation of graphite, just as a
solution when slowly cooled yields a portion of the substance
dissolved. The separation of carbon in the form of graphite on the
conversion of white cast iron into grey is never complete, however
slowly the separation be carried on; part of the carbon remains in
combination with the iron in the same state in which it exists in
white cast iron. Hence when grey cast iron is treated with acids,
the whole of the carbon does not remain in the form of graphite,
but a part of it is separated as hydrocarbons, which proves the
existence of chemically-combined carbon in grey cast iron. It is
sufficient to re-melt grey cast iron and to cool it quickly to
transform it into white cast iron. It is not carbon alone that
influences the properties of cast iron; when it contains a
considerable amount of sulphur, cast iron remains white even after
having been slowly cooled. The same is observed in cast iron very
rich in manganese (5 to 7 p.c.), and in this latter case the
fracture is very distinctly crystalline and brilliant. When cast
iron contains a large amount of manganese, the quantity of carbon
may also be increased. Crystalline varieties of cast iron rich in
manganese are in practice called ferro-manganese (p. 310), and are
prepared for the Bessemer process. Grey cast iron not having an
uniform structure is much more liable to various changes than dense
and thoroughly uniform white cast iron, and the latter oxidises
much more slowly in air than the former. White cast iron is not
only used for conversion into wrought iron and steel, but also in
those cases where great hardness is required, although it be
accompanied by a certain brittleness; for instance, for making
rollers, plough-shares, &c.
Cast iron is a material which is either suitable for direct
application for casting in moulds or else for working up into _wrought
iron_ and _steel_. The latter principally differ from cast iron in their
containing less carbon--thus, steel contains from 1 p.c. to 0·5 p.c. of
carbon and far less silicon and manganese than cast iron; wrought iron
does not generally contain more than 0·25 p.c. of carbon and not more
than 0·25 p.c. of the other impurities. Thus the essence of the working
up of cast iron into steel and wrought iron consists in the removal of
the greater part of the carbon and other elements, S, P, Mn, Si, &c. This
is effected by means of oxidation, because the oxygen of the atmosphere,
oxidising the iron at a high temperature, forms solid oxides with it; and
the latter, coming into contact with the carbon contained in the cast
iron, are deoxidised, forming wrought iron and carbonic oxide, which is
evolved from the mass in a gaseous form. It is evident that the oxidation
must be carried on with a molten mass in a state of agitation, so that
the oxygen of the air may be brought into contact with the whole mass of
carbon contained in the cast iron, or else the operation is effected by
means of the addition of oxygen compounds of iron (oxides, ores, as in
Martin's process). Cast iron melts much more easily than wrought iron and
steel, and, therefore, as the carbon separates, the mass in the furnace
(in puddling) or hearth (in the bloomery process) becomes more and more
solid; moreover the degree of hardness forms, to a certain extent, a
measure of the amount of carbon separated, and the operation may
terminate either in the formation of steel or wrought iron.[10] In any
case, the iron used for industrial purposes contains impurities.
_Chemically pure iron_ may be obtained by precipitating iron from a
solution (a mixture of ferrous sulphate with magnesium sulphate or
ammonium chloride) by the prolonged action of a feeble galvanic current;
the iron may be then obtained as a dense mass. This method, proposed by
Böttcher and applied by Klein, gives, as R. Lenz showed, iron containing
occluded hydrogen, which is disengaged on heating. This galvanic
deposition of iron is used for making galvanoplastic _clichés_, which are
distinguished for their great hardness. Electro-deposited iron is
brittle, but if heated (after the separation of the hydrogen) it becomes
soft. If pure ferric hydroxide, which is easily prepared by the
precipitation of solutions of ferric salts by means of ammonia, be heated
in a stream of hydrogen, it forms, first of all, a dull black powder
which ignites spontaneously in air (pyrophoric iron), and then a grey
powder of pure iron. The powdery substance first obtained is an iron
suboxide; when thrown into the air it ignites, forming the oxide
Fe_{3}O_{4}. If the heating in hydrogen be continued, more water and pure
iron, which does not ignite spontaneously, will be obtained. If a small
quantity of iron be fused in the oxyhydrogen flame (with an excess of
oxygen) in a piece of lime and mixed with powdered glass, pure molten
iron will be formed, because in the oxyhydrogen flame iron melts and
burns, but the substances mixed with the iron oxidise first. The oxidised
impurities here either disappear (carbonic anhydride) in a gaseous form,
or turn into slag (silica, manganese, oxide, and others)--that is, fuse
with the glass. Pure iron has a silvery white colour and a specific
gravity of 7·84; it melts at a temperature higher than the melting-points
of silver, gold, nickel, and steel, _i.e._ about 1400°-1500° and below
the melting point of platinum (1750°).[11] But pure iron becomes soft at
a temperature considerably below that at which it melts, and may then be
easily forged, welded, and rolled or drawn into sheets and wire.[11 bis]
Pure iron may be rolled into an exceedingly thin sheet, weighing less
than a sheet of ordinary paper of the same size. This ductility is the
most important property of iron in all its forms, and is most marked with
sheet iron, and least so with cast iron, whose ductility, compared with
wrought iron, is small, but it is still very considerable when compared
with other substances--such, for instance, as rocks.[12]
[10] This direct process of separating the carbon from cast iron is
termed _puddling_. It is conducted in reverberatory furnaces. The
cast iron is placed on the bed of the furnace and melted; through
a special aperture, the puddler stirs up the oxidising mass of
cast iron, pressing the oxides into the molten iron. This
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]