strong hydrochloric acid and a small quantity of water, and gently
warmed, it all passes into solution, and no chlorine is evolved; on
cooling, the liquid deposits red prismatic crystals, known as
_Peligot's salt_, very stable in air. This has the composition
KCl,CrO_{3}, and is formed according to the equation
K_{2}Cr_{2}O_{7} + 2HCl = 2KCl,CrO_{3} + H_{2}O. It is evident that
this is the first chloranhydride of chromic acid, HCrO_{3}Cl, in
which the hydrogen is replaced by potassium. It is decomposed by
water, and on evaporation the solution yields potassium dichromate
and hydrochloric acid. This is a fresh instance of the reversible
reactions so frequently encountered. With sulphuric acid Peligot's
salt forms chromyl chloride. The latter circumstance, and the fact
that Geuther produced Peligot's salt from potassium chromate and
chromyl chloride, give reason for thinking that it is a compound of
these two substances: 2KCl,CrO_{3} = K_{2}CrO_{4} + CrO_{2}Cl_{2}.
It is also sometimes regarded as potassium dichromate in which one
atom of oxygen is replaced by chlorine--that is,
K_{2}Cr_{2}O_{6}Cl_{2}, corresponding with K_{2}Cr_{2}O_{7}. When
heated it parts with all its chlorine, and on further heating gives
chromic oxide.
[6] This intermediate degree of oxidation, CrO_{2}, may also be
obtained by mixing solutions of chromic salts with solutions of
chromates. The brown precipitate formed contains a compound,
Cr_{2}O_{3},CrO_{3}, consisting of equivalent amounts of chromic
oxide and anhydride. The brown precipitate of chromium dioxide
contains water. The same substance is formed by the imperfect
deoxidation of chromic anhydride by various reducing agents.
Chromic oxide, when heated, absorbs oxygen, and appears to give the
same substance. Chromic nitrate, when ignited, also gives this
substance. When this substance is heated it first disengages water
and then oxygen, chromic oxide being left. It corresponds with
manganese dioxide, Cr_{2}O_{3},CrO_{3} = 3CrO_{2}. Krüger treated
chromium dioxide with a mixture of sodium chloride and sulphuric
acid, and found that chlorine gas was evolved, but that chromyl
chloride was not formed. Under the action of light, a solution of
chromic acid also deposits the brown dioxide. At the ordinary
temperature chromic anhydride leaves a brown stain upon the skin
and tissues, which probably proceeds from a decomposition of the
same kind. Chromic anhydride is soluble in alcohol containing
water, and this solution is decomposed in a similar manner by
light. Chromium dioxide forms K_{2}CrO_{4} when treated with
H_{2}O_{2} in the presence of KHO.
The action of hydrogen peroxide on a solution of chromic acid or of
potassium dichromate gives a blue solution, which very quickly becomes
colourless with the disengagement of oxygen. Barreswill showed that this
is due to the formation of a _perchromic anhydride_, Cr_{2}O_{7},
corresponding with sulphur peroxide. This peroxide is remarkable from the
fact that it very easily dissolves in ether and is much more stable in
this solution, so that, by shaking up hydrogen peroxide mixed with a
small quantity of chromic acid, with ether, it is possible to transfer
all the blue substance formed to the ether.[6 bis]
[6 bis] Now that persulphuric acid H_{2}S_{2}O_{8} is well known it
might be supposed that perchromic anhydride, Cr_{2}O_{7}, would
correspond to perchromic acid, H_{2}Cr_{2}O_{8}, but as yet it is
not certain whether corresponding salts are formed. Péchard (1891)
on adding an excess of H_{2}O_{2} and baryta water to a dilute
solution of CrO_{2} (8 grm. per litre), observed the formation of a
yellow precipitate, but oxygen was disengaged at the same time and
the precipitate (which easily exploded when dried) was found to
contain, besides an admixture of BaO_{2}, a compound BaCrO_{5}, and
this = BaO_{2} + CrO_{3}, and does not correspond to perchromic
acid. The fact of its decomposing with an explosion, and the mode
of its preparation, proves, however, that this is a similar
derivative of peroxide of hydrogen to persulphuric acid (Chapter
XX.)
With oxygen acids, chromic acid evolves oxygen; for example, with
sulphuric acid the following reaction takes place: 2CrO_{3} +
3H_{2}SO_{4} = Cr_{2}(SO_{4})_{3} + O_{3} + 3H_{2}O. It will be readily
understood from this that _a mixture of chromic acid_ or _of its salts
with sulphuric acid_ forms an excellent _oxidising agent_, which is
frequently employed in chemical laboratories and even for technical
purposes as a means of oxidation. Thus hydrogen sulphide and sulphurous
anhydride are converted into sulphuric acid by this means. Chromic acid
is able to act as a powerful oxidising agent because it passes into
chromic oxide, and in so doing disengages half of the oxygen contained in
it: 2CrO_{3} = Cr_{2}O_{3} + O_{3}. Thus chromic anhydride itself is a
powerful oxidising agent, and is therefore employed instead of nitric
acid in galvanic batteries (as a depolariser), the hydrogen evolved at
the carbon being then oxidised, and the chromic acid converted into a
non-volatile product of deoxidation, instead of yielding, as nitric acid
does, volatile lower oxides of offensive odour. Organic substances are
more or less perfectly oxidised by means of chromic anhydride, although
this generally requires the aid of heat, and does not proceed in the
presence of alkalis, but generally _in the presence of acids_. In acting
on a solution of potassium iodide, chromic acid, like many oxidising
agents, liberates iodine; the reaction proceeds in proportion to the
amount of CrO_{3} present, and may serve for determining the amount of
CrO_{3}, since the amount of iodine liberated can be accurately
determined by the iodometric method (Chapter XX., Note 42). If chromic
anhydride be ignited in a stream of ammonia, it gives chromic oxide,
water, and nitrogen. In all cases when chromic acid acts as an oxidising
agent in the presence of acids and under the action of heat, the product
of its deoxidation is a chromic salt, CrX_{3}, which is characterised by
the green colour of its solution, so that the _red_ or yellow _solution_
of a salt of chromic acid is then transformed into a _green solution_ of
a chromic salt, derived from chromic oxide, Cr_{2}O_{3}, which is closely
analogous to Al_{2}O_{3}, Fe_{2}O_{3}, and other bases of the composition
R_{2}O_{3}. This analogy is seen in the insolubility of the anhydrous
oxide, in the gelatinous form of the colloidal hydrate, in the formation
of alums,[7] of a volatile chloride of chromium, &c.[7 bis]
[7] As a mixture of potassium dichromate and sulphuric acid is usually
employed for oxidation, the resultant solution generally contains a
double sulphate of potassium and chromium--that is, _chrome alum_,
isomorphous with ordinary alum--K_{2}Cr_{2}O_{7} + 4H_{2}SO_{4} +
20H_{2}O = O_{3} + K_{2}Cr_{2}(SO_{4})_{4},24H_{2}O or
2(KCr(SO_{4})_{2},12H_{2}O). It is prepared by dissolving potassium
dichromate in dilute sulphuric acid; alcohol is then added and the
solution slightly heated, or sulphurous anhydride is passed through
it. On the addition of alcohol to a cold mixture of potassium
dichromate and sulphuric acid, the gradual disengagement of
pleasant-smelling volatile products of the oxidation of alcohol,
and especially of aldehyde, C_{2}H_{4}O, is remarked. If the
temperature of decomposition does not exceed 35°, a _violet_
solution of chrome alum is obtained, but if the temperature be
higher, a solution of the same alum is obtained of a _green_
colour. As chrome alum requires for solution 7 parts of water at
the ordinary temperature, it follows that if a somewhat strong
solution of potassium dichromate be taken (4 parts of water and
1-1/2 of sulphuric acid to 1 part of dichromate), it will give so
concentrated a solution of chrome alum that on cooling, the salt
will separate without further evaporation. _If the liquid_,
prepared as above or in any instance of the deoxidation of chromic
acid, _be heated_ (the oxidation naturally proceeds more rapidly)
somewhat strongly, for instance, to the boiling-point of water, or
if the violet solution already formed be raised to the same
temperature, it acquires a bright _green colour_, and on
evaporation the same mixture, which at lower temperatures so easily
gives cubical crystals of chrome alum, _does not give any crystals
whatever_. _If the green solution be kept_, however, _for several
weeks_ at the ordinary temperature, it deposits _violet crystals_
of chrome alum. The green solution, when evaporated, gives a
non-crystalline mass, and the violet crystals lose water at 100°
and turn green. It must be remarked that the transition of the
green modification into the violet is accompanied by a decrease in
volume (Lecoq de Boisbaudran, Favre). If the green mass formed at
the higher temperature be evaporated to dryness and heated at 30°
in a current of air, it does not retain more then 6 equivalents of
water. Hence Löwel, and also Schrötter, concluded that the green
and violet modifications of the alum depend on different degrees of
combination with water, which may be likened to the different
compounds of sodium sulphate with water and to the different
hydrates of ferric oxide.
However, the question in this case is not so simple, as we shall
afterwards see. Not chrome alum alone, but _all the chromic salts_,
give two, if not three, _varieties_. At least, there is no doubt
about the existence of two--a _green_ and a _violet modification_.
The green chromic salts are obtained by heating solutions of the
violet salts, the violet solutions are produced on keeping
solutions of the green salts for a long time. The conversion of the
violet salts into green by the action of heat is itself an
indication of the possibility of explaining the different
modifications by their containing different proportions (or states)
of water, and, moreover, by the green salts having a less amount of
water than the violet. However, there are other explanations.
Chromic oxide is a base like alumina, and is therefore able to give
both acid and basic salts. It is supposed that the difference
between the green and violet salts is due to this fact. This
opinion of Krüger is based on the fact that alcohol separates out a
salt from the green solution which contains less sulphuric acid
than the normal violet salt. On the other hand, Löwel showed that
all the acid cannot be separated from the green chromic salts by
suitable reagents, as easily as it can be from the same solution of
the violet salts; thus barium salts do not precipitate all the
sulphuric acid from solutions of the green salts. According to
other researches the cause of the varieties of the chromic salts
lies in a difference in the bases they contain--that is, it is
connected with a modification of the properties of the oxide of
chromium itself. This only refers to the hydroxides, but as
hydroxides themselves are only special forms of salts, the
differences observed as yet in this direction between the
hydroxides only confirm the generality of the difference observed
in the chromic compounds (_see_ Note 7 bis).
The salts of chromic oxide, like those of alumina, are easily
decomposed, give basic and double salts, and have an acid reaction,
as chromic oxide is a feeble base. Potassium and sodium hydroxides
give a _precipitate_ of the hydroxide with chromic salts, CrX_{3}.
The violet and green salts give a _hydroxide soluble in an excess
of the reagent_; but the hydroxide is held in solution by very
feeble affinities, so that it is partially separated by heat and
dilution with water, and completely so on boiling. In an alkaline
solution, chromic hydroxide is easily converted into chromic acid
by the action of lead dioxide, chlorine, and other oxidising
agents. If the chromic oxide occurs together with such oxides as
magnesia, or zinc oxide, then on precipitation it separates out
from its solution in combination with these oxides, forming, for
example, ZnO,Cr_{2}O_{3}. Viard obtained compounds of Cr_{2}O_{3}
with the oxides of Mg, Zn, Cd, &c.) On precipitating the violet
solution of chrome alum with ammonia, a precipitate containing
Cr_{2}O_{3},6H_{2}O is obtained, whilst the precipitate from the
boiling solution with caustic potash was a hydrate containing four
equivalents of water. When fused with borax chromic salts give a
green glass. The same coloration is communicated to ordinary glass
by the presence of traces of chromic oxide. A chrome glass
containing a large amount of chromic oxide may be ground up and
used as a green pigment. Among the hydrates of oxide of chromium
_Guignet's green_ forms one of the widely-used green pigments which
have been substituted for the poisonous arsenical copper pigments,
such as Schweinfurt green, which formerly was much used. Guignet's
green has an extremely bright green colour, and is distinguished
for its great stability, not only under the action of light but
also towards reagents; thus it is not altered by alkaline
solutions, and even nitric acid does not act on it. This pigment
remains unchanged up to a temperature of 250°; it contains
Cr_{2}O_{3},2H_{2}O, and generally a small amount of alkali. It is
prepared by fusing 3 parts of boric acid with 1 part of potassium
dichromate; oxygen is disengaged, and a green glass, containing a
mixture of the borates of chromium and potassium, is obtained. When
cool this glass is ground up and treated with water, which extracts
the boric acid and alkali and leaves the above-named chromic
hydroxide behind. This hydroxide only parts with its water at a red
heat, leaving the anhydrous oxide.
The chromic hydroxides lose their water by ignition, and in so
doing become spontaneously incandescent, like the ordinary ferric
hydroxide (Chapter XXII.). It is not known, however, whether all
the modifications of chromic oxide show this phenomenon. The
anhydrous _chromic oxide_, Cr_{2}O_{3}, is exceedingly difficultly
soluble in acids, if it has passed through the above recalescence.
But if it has parted with its water, or the greater part of it, and
not yet undergone this self-induced incandescence (has not lost a
portion of its energy), then it is soluble in acids. It is not
reduced by hydrogen. It is easily obtained in various crystalline
forms by many methods. The chromates of mercury and ammonium give a
very convenient method for its preparation, because when ignited
they leave chromic oxide behind. In the first instance oxygen and
mercury are disengaged, and in the second case nitrogen and water:
2Hg_{2}CrO_{4} = Cr_{2}O_{3} + O_{5} + 4Hg or
(NH_{4})_{2}Cr_{2}O_{7} = Cr_{2}O_{3} + 4H_{2}O + N_{2}. The second
reaction is very energetic, and the mass of salt burns
spontaneously if the temperature be sufficiently high. A mixture of
potassium sulphate and chromic oxide is formed by heating potassium
dichromate with an equal weight of sulphur: K_{2}Cr_{2}O_{7} + S =
K_{2}SO_{4} + Cr_{2}O_{3}. The sulphate is easily extracted by
water, and there remains a bright green residue of the oxide, whose
colour is more brilliant the lower the temperature of the
decomposition. The oxide thus obtained is used as a green pigment
for china and enamel. The anhydrous chromic oxide obtained from
chromyl chloride, CrO_{2}Cl_{2}, has a specific gravity of 5·21,
and forms almost black crystals, which give a green powder. They
are hard enough to scratch glass, and have a metallic lustre. The
crystalline form of chromic oxide is identical with that of the
oxide of iron and alumina, with which it is isomorphous.
[7 bis] The most important of the compounds corresponding with chromic
oxide is _chromic chloride_, Cr_{2}Cl_{6}, which is known in an
anhydrous and in a hydrated form. It resembles ferric and aluminic
chlorides in many respects. There is a great difference between the
anhydrous and the hydrated chlorides; the former is insoluble in
water, the latter easily dissolves, and on evaporation its solution
forms a hygroscopic mass which is very unstable and easily evolves
hydrochloric acid when heated with water. The anhydrous form is of
a violet colour, and Wöhler gives the following method for its
preparation: an intimate mixture is prepared of the anhydrous
chromic oxide with carbon and organic matter, and charged into a
wide infusible glass or porcelain tube which is heated in a
combustion furnace; one extremity of the tube communicates with an
apparatus generating chlorine which is passed through several
bottles containing sulphuric acid in order to dry it perfectly
before it reaches the tube. On heating the portion of the tube in
which the mixture is placed and passing chlorine through, a
slightly volatile sublimate of chromic chloride, CrCl_{3} or
Cr_{2}Cl_{6}, is formed. This substance forms _violet tabular
crystals_, which may be distilled in dry chlorine without change,
but which, however, require a red heat for their volatilisation.
These crystals are greasy to the touch and insoluble in water, but
if they be powdered and boiled in water for a long time they pass
into _a green solution_. Strong sulphuric acid does not act on the
anhydrous salt, or only acts with exceeding slowness, like water.
Even aqua regia and other acids do not act on the crystals, and
alkalis only show a very feeble action. The specific gravity of the
crystals is 2·99. When fused with sodium carbonate and nitre they
give sodium chloride and potassium chromate, and when ignited in
air they form green chromic oxide and evolve chlorine. On ignition
in a stream of ammonia, chromic chloride forms sal-ammoniac and
chromium nitride, CrN (analogous to the nitrides BN, AlN). Mosberg
and Peligot showed that when chromic chloride is ignited in
hydrogen, it parts with one-third of its chlorine, forming chromous
chloride, CrCl_{2}--that is, there is formed from a compound
corresponding with chromic oxide, Cr_{2}O_{3}, a compound answering
to the _suboxide_, chromous oxide, CrO--just as hydrogen converts
ferric chloride into ferrous chloride with the aid of heat.
_Chromous chloride_, CrCl_{2}, forms colourless crystals easily
soluble in water, which in dissolving evolve a considerable amount
of heat, and form a blue liquid, capable of absorbing oxygen from
the air with great facility, being converted thereby into a chromic
compound.
The blue solution of chromous chloride may also be obtained by the
action of metallic zinc on the green solution of the hydrated
chromic chloride; the zinc in this case takes up chlorine just as
the hydrogen did. It must be employed in a large excess. Chromic
oxide is also formed in the action of zinc on chromic chloride, and
if the solution remain for a long time in contact with the zinc the
whole of the chromium is converted into chromic oxychloride. Other
chromic salts are also reduced by zinc into _chromous salts_,
CrX_{2}, just as the ferric salts FeX_{3} are converted into
ferrous salts FeX_{2} by it. The chromous salts are exceedingly
unstable and easily oxidise and pass into chromic salts; hence the
reducing power of these salts is very great. From cupric salts they
separate cuprous salts, from stannous salts they precipitate
metallic tin, they reduce mercuric salts into mercurous and ferric
into ferrous salts. Moreover, they absorb oxygen from the air
directly. With potassium chromate they give a brown precipitate of
chromium dioxide or of chromic oxide, according to the relative
amounts of the substances taken: CrO_{3} + CrO = 2CrO_{2} or
CrO_{3} + 3CrO = 2Cr_{2}O_{3}. Aqueous ammonia gives a blue
precipitate, and in the presence of ammoniacal salts a blue liquid
is obtained which turns red in the air from oxidation. This is
accompanied by the formation of compounds analogous to those given
by cobalt (Chapter XXII.). A solution of chromous chloride with a
hot saturated solution of sodium acetate, C_{2}H_{3}NaO_{2}, gives,
on cooling, transparent red crystals of chromous acetate,
C_{4}H_{6}CrO_{4},H_{2}O. This salt is also a powerful reducing
agent, but may be kept for a long time in a vessel full of carbonic
anhydride.
The insoluble anhydrous _chromic chloride_ CrCl_{3} very easily
_passes into solution_ in the presence of a trace (0·004) of
_chromous chloride_ CrCl_{2}. This remarkable phenomenon was
observed by Peligot and explained by Löwel in the following manner:
chromous chloride, as a lower stage of oxidation, is capable of
absorbing both oxygen and chlorine, combining with various
substances. It is able to decompose many chlorides by taking up
chlorine from them; thus it precipitates mercurous chloride from a
solution of mercuric chloride, and in so doing passes into chromic
chloride: 2CrCl_{2} + 2HgCl_{2} = Cr_{2}Cl_{6} + 2HgCl. Let us
suppose that the same phenomenon takes place when the anhydrous
chromic chloride is mixed with a solution of chromous chloride. The
latter will then take up a portion of the chlorine of the former,
and pass into a soluble hydrate of chromic chloride (hydrochloride
of oxide of chromium), and the original anhydrous chromic chloride
will pass into chromous chloride. The chromous chloride re-formed
in this manner will then act on a fresh quantity of the chromic
chloride, and in this manner transfer it entirely into solution as
hydrate. This view is confirmed by the fact that other chlorides,
capable of absorbing chlorine like chromous chloride, also induce
the solution of the insoluble chromic chloride--for example,
ferrous chloride, FeCl_{2}, and cuprous chloride. The presence of
zinc also aids the solution of chromic chloride, owing to its
converting a portion of it into chromous chloride. The solution of
chromic chloride in water obtained by these methods is perfectly
identical with that which is formed by dissolving chromic hydroxide
in hydrochloric acid. On evaporating the _green solution_ obtained
in this manner, it gives a green mass, containing water. On further
heating it leaves a soluble chromic oxychloride, and when ignited
it first forms an insoluble oxychloride and then chromic oxide; but
no anhydrous chromic chloride, Cr_{2}Cl_{6}, is formed by heating
the aqueous solution of chromic chloride, which forms an important
fact in support of the view that the green solution of chromic
chloride is nothing else but hydrochloride of oxide of chromium. At
100° the composition of the green hydrate is Cr_{2}Cl_{6},9H_{2}O,
and on evaporation at the ordinary temperature over H_{2}SO_{4}
crystals are obtained with 12 equivalents of water; the red mass
obtained at 120° contains Cr_{2}O_{3},4Cr_{2}Cl_{6},24H_{2}O. The
greater portion of it is soluble in water, like the mass which is
formed at 150°. The latter contains
Cr_{2}O_{3},2Cr_{2}Cl_{6},9H_{2}O = 3(Cr_{2}OCl_{4},3H_{2}O)--that
is, it presents the same composition as chromic chloride in which
one atom of oxygen replaces two of chlorine. And if the hydrate of
chromic chloride be regarded as Cr_{2}O_{3},6HCl, the substance
which is obtained should be regarded as Cr_{2}O_{3},4HCl combined
with water, H_{2}O. The addition of alkalis--for example,
baryta--to a solution of chromic chloride immediately produces a
precipitate, which, however, re-dissolves on shaking, owing to the
formation of one of the oxychlorides just mentioned, which may be
regarded as _basic salts_. Thus we may represent the product of the
change produced on chromic chloride under the influence of water
and heat by the following formulæ: first Cr_{2}O_{3},6HCl or
Cr_{2}Cl_{6},3H_{2}O is formed, then Cr_{2}O_{3},4HCl,H_{2}O or
Cr_{2}OCl_{4},3H_{2}O, and lastly Cr_{2}O_{3},2HCl,2H_{2}O or
Cr_{2}O_{2}Cl_{2},3H_{2}O. In all three cases there are 2
equivalents of chromium to at least 3 equivalents of water. These
compounds may be regarded as being intermediate between chromic
hydroxide and chloride; chromic chloride is Cr_{2}Cl_{6}, the first
oxychloride Cr_{2}(OH)_{2}Cl_{4}, the second Cr_{2}(OH)_{4}Cl_{2},
and the hydrate Cr_{2}(OH)_{6}--that is, the chlorine is replaced
by hydroxyl.
It is very important to remark two circumstances in respect to
this: (1) That the whole of the chlorine in the above compounds is
not precipitated from their solutions by silver nitrate; thus the
normal salt of the composition Cr_{2}Cl_{6},9H_{2}O only gives up
two-thirds of its chlorine; therefore Peligot supposes that the
normal salt contains the oxychloride combined with hydrochloric
acid: Cr_{2}Cl_{6} + 2H_{2}O = Cr_{2}O_{2}Cl_{2},4HCl, and that the
chlorine held as hydrochloric acid reacts with the silver, whilst
that held in the oxychloride does not enter into reaction, just as
we observe a very feebly-developed faculty for reaction in the
anhydrous chromic chloride; and (2) if the green aqueous solution
of CrCl_{3} be left to stand for some time, it ultimately turns
violet; in this form the whole of the chlorine is precipitated by
AgNO_{3}, whilst boiling re-converts it into the green variety.
Löwel obtained the violet solution of hydrochloride of chromic
oxide by decomposing the violet chromic sulphate with barium
chloride. Silver nitrate precipitates all the chlorine from this
violet modification; but if the violet solution be boiled and so
converted into the green modification, silver nitrate then only
precipitates a portion of the chlorine.
Recoura (1890-1893) obtained a crystallohydrate of violet chromium
sulphate, Cr_{2}(SO_{4})_{3}, with 18 or 15 H_{2}O. By boiling a
solution of this crystallohydrate, he converted it into the green
salt, which, when treated with alkalis, gave a precipitate of
Cr_{2}O_{3},2H_{2}O, soluble in 2H_{2}SO_{4} (and not 3), and only
forming the basic salt, Cr_{2}(OH)_{2}(SO_{4})_{2}. He therefore
concludes that the green salts are basic salts. The cryoscopic
determinations made by A. Speransky (1892) and Marchetti (1892)
give a greater 'depression' for the violet than the green salts,
that is, indicate a greater molecular weight for the green salts.
But as Étard, by heating the violet sulphate to 100°, converted it
into a green salt of the same composition, but with a smaller
amount of H_{2}O, it follows that the formation of a basic salt
alone is insufficient to explain the difference between the green
and violet varieties, and this is also shown by the fact that
BaCl_{2} precipitates the whole of the sulphuric acid of the violet
salt, and only a portion of that of the green salt. A. Speransky
also showed that the molecular electro-conductivity of the green
solutions is less than that of the violet. It is also known that
the passage of the former into the latter is accompanied by an
increase of volume, and, according to Recoura, by an evolution of
heat also.
Piccini's researches (1894) throw an important light upon the
peculiarities of the green chromium trichloride (or chromic
chloride); he showed (1) that AgF (in contradistinction to the
other salts of silver) precipitates all the chlorine from an
aqueous solution of the green variety; (2) that solutions of green
CrCl_{3},6H_{2}O in ethyl alcohol and acetone precipitate all their
chlorine when mixed with a similar solution of AgNO_{3}; (3) that
the rise of the boiling-point of the ethyl alcohol and acetone
green solutions of CrCl_{3},6H_{2}O (Chapter VII., Note 27 bis)
shows that i in this case (as in the aqueous solutions of MgSO_{4}
and HgCl_{2}) is nearly equal to 1, that is, that they are like
solutions of non-conductors; (4) that a solution of green CrCl_{3}
in methyl alcohol at first precipitates about 7/8 of its chlorine
(an aqueous solution about 2/3) when treated with AgNO_{3}, but
after a time the whole of the chlorine is precipitated; and (5)
that an aqueous solution of the green variety gradually passes into
the violet, while a methyl alcoholic solution preserves its green
colour, both of itself and also after the whole of the chlorine has
been precipitated by AgNO_{3}. If, however, in an aqueous or methyl
alcoholic solution only a portion of the chlorine be precipitated,
the solution gradually turns violet. In my opinion the general
meaning of all these observations requires further elucidation and
explanation, which should be in harmony with the theory of
solutions. Recoura, moreover, obtained compounds of the green salt,
Cr_{2}(SO_{4})_{3}, with 1, 2, and 3 molecules of H_{2}SO_{4},
K_{2}SO_{4}, and even a compound Cr_{2}(SO_{4})_{3}H_{2}CrO_{4}. By
neutralising the sulphuric acid of the compounds of
Cr_{2}(SO_{4})_{3} and H_{2}SO_{4} with caustic soda, Recoura
obtained an evolution of 33 thousand calories per each 2NaHO, while
free H_{2}SO_{4} only gives 30·8 thousand calories. Recoura is of
opinion that special _chromo sulphuric acids_, for instance
(CrSO_{4})H_{2}SO_{4} = 1/2Cr_{2}(SO_{4})_{3}H_{2}SO_{4}, are
formed. With a still larger excess of sulphuric acid, Recoura
obtained salts containing a still greater number of sulphuric acid
radicles, but even this method does not explain the difference
between the green and violet salts.
These facts must naturally be taken into consideration in order to
arrive at any complete decision as to the cause of the different
modifications of the chromic salts. We may observe that the green
modification of chromic chloride does not give double salts with
the metallic chlorides, whilst the violet variety forms compounds
Cr_{2}Cl_{6},2RCl (where R = an alkali metal), which are obtained
by heating the chromates with an excess of hydrochloric acid and
evaporating the solution until it acquires a violet colour. As the
result of all the existing researches on the green and violet
chromic salts, it appears to me most probable that their difference
is determined by the feeble basic character of chromic oxide, by
its faculty of giving basic salts, and by the colloidal properties
of its hydroxide (these three properties are mutually connected),
and moreover, it seems to me that the relation between the green
and violet salts of chromic oxide best answers to the relation of
the purpureo to the luteo cobaltic salts (Chapter XXII., Note 35).
This subject cannot yet be considered as exhausted (_see_ Note 7).
We may here observe that with tin the chromic salts, CrX_{3}, give
at low temperatures CrX_{2} and SnX_{2}, whilst at high
temperatures, on the contrary, CrX_{2} reduces the metal from its
salts SnX_{2}. The reaction, therefore, belongs to the class of
reversible reactions (Beketoff).
Poulenc obtained anhydrous CrF_{3} (sp. gr. 3·78) and CrF_{2} (sp.
gr. 4·11) by the action of gaseous HF upon CrCl_{2}. A solution of
fluoride of chromium is employed as a mordant in dyeing. Recoura
(1890) obtained green and violet varieties of Cr_{2}Br_{6},6H_{2}O.
The green variety can only be kept in the presence of an excess of
HBr in the solution; if alone its solution easily passes into the
violet variety with evolution of heat.
_Chromic oxide_, Cr_{2}O_{3}, rarely found, and in small quantities,
in chrome ochre, is formed by the oxidation of chromium and its lower
oxides, by the reduction of chromates (for example, of ammonium or
mercuric chromate) and by the decomposition (splitting up) of the saline
compounds of the oxide itself, CrX_{3} or Cr_{2}X_{6}, like alumina,
which it resembles in forming a feeble base easily giving double and
basic salts, which are either green or violet.
The reduction of chromic oxide--for instance, in a solution by zinc
and sulphuric acid--leads to the formation of chromous oxide, CrO, and
its salts, CrX_{2}, of a blue colour (_see_ Notes 7 and 7 ^{bis}). The
further reduction[8] of oxide of chromium and its corresponding compounds
gives _metallic chromium_. Deville obtained it (probably containing
carbon) by reducing chromic oxide with carbon, at a temperature near the
melting point of platinum, about 1750°, but the metal itself does not
fuse at this temperature. Chromium has a steel-grey colour and is very
hard (sp. gr. 5·9), takes a good polish, and dissolves in hydrochloric
acid, but cold dilute sulphuric and nitric acids have no action upon it.
Bunsen obtained metallic chromium by decomposing a solution of chromic
chloride, Cr_{2}Cl_{6}, by a galvanic current, as scales of a grey colour
(sp. gr. 7·3). Wöhler obtained crystalline chromium by igniting a mixture
of the anhydrous chromic chloride Cr_{2}Cl_{6} (_see_ Note 7 bis) with
finely-divided zinc, and sodium and potassium chlorides, at the
boiling-point of zinc. When the resultant mass has cooled the zinc may be
dissolved in dilute nitric acid, and grey crystalline chromium (sp. gr.
6·81) is left behind. Frémy also prepared crystalline chromium by the
action of the vapour of sodium on anhydrous chromic chloride in a stream
of hydrogen, using the apparatus shown in the accompanying drawing, and
placing the sodium and the chromic chloride in separate porcelain boats.
The tube containing these boats is only heated when it is quite full of
dry hydrogen. The crystals of metallic chromium obtained in the tube are
grey cubes having a considerable hardness and withstanding the action of
powerful acids, and even of aqua regia. The chromium obtained by Wöhler
by the action of a galvanic current is, on the contrary, acted on under
these conditions. The reason of this difference must be looked for in the
presence of impurities, and in the crystalline structure. But in any
case, among the properties of metallic chromium, the following may be
considered established: it is white in colour, with a specific gravity of
about 6·7, is extremely hard in a crystalline form, is not oxidised by
air at the ordinary temperature, and with carbon it forms alloys like
cast iron and steel.
[8] The reduction of metallic chromium proceeds with comparative ease
in aqueous solutions. Thus the action of sodium amalgams upon a
strong solution of Cr_{2}Cl_{6} gives (first CrCl_{2}) an amalgam
of chromium from which the mercury may be easily driven off by
heating (in hydrogen to avoid oxidation), and there remains a
spongy mass of easily oxidizable chromium. Plaset (1891), by
passing an electric current through a solution of chrome alum mixed
with a small amount of H_{2}SO_{4} and K_{2}SO_{4}, obtained hard
scales of chromium of a bluish-white colour possessing great
hardness and stability (under the action of water, air, and acids).
Glatzel (1890) reduced a mixture of 2KCl + Cr_{2}Cl_{6} by heating
it to redness with shavings of magnesium. The metallic chromium
thus obtained has the appearance of a fine light-grey powder which
is seen to be crystalline under the microscope; its sp. gr. at 16°
is 6·7284. It fuses (with anhydrous borax) only at the highest
temperatures, and after fusion presents a silver-white fracture.
The strongest magnet has no action upon it.
Moissan (1893) obtained chromium by reducing the oxide Cr_{2}O_{3}
with carbon in the electrical furnace (Chapter VIII., Note 17) in
9-10 minutes with a current of 350 ampères and 50 volts. The
mixture of oxide and carbon gives a bright ingot weighing 100-110
grams. A current of 100 ampères and 50 volts completes the
experiment upon a smaller quantity of material in 15 minutes; a
current of 30 ampères and 50 volts gave an ingot of 10 grams in
30-40 minutes. The resultant carbon alloy is more or less rich in
chromium (from 87·37-91·7 p.c.). To obtain the metal free from
carbon, the alloy is broken into large lumps, mixed with oxide of
chromium, put into a crucible and covered with a layer of oxide.
This mixture is then heated in the electric furnace and the pure
metal is obtained. This reduction can also be carried on with
chrome iron ore FeOCr_{2}O_{3} which occurs in nature. In this case
a homogeneous alloy of iron and chromium is obtained. If this alloy
be thrown in lumps into molten nitre, it forms insoluble
sesquioxide of iron and a soluble alkaline chromate. This alloy of
iron and chromium dissolved in molten steel (chrome steel) renders
it hard and tough, so that such steel has many valuable
applications. The alloy, containing about 3 p.c. Cr and about 1·3
p.c. carbon, is even harder than the ordinary kinds of tempered
steel and has a fine granular fracture. The usual mode of preparing
the ferrochromes for adding to steel is by fusing powdered chrome
iron ore under fluxes in a graphite crucible.
[Illustration: FIG. 92.--Apparatus for the preparation of metallic
chromium by igniting chromic chloride and sodium in a stream of
hydrogen.]
The two analogues of chromium, _molybdenum_ and _tungsten_ (or wolfram),
are of still rarer occurrence in nature, and form acid oxides, RO_{3},
which are still less energetic than CrO_{3}. Tungsten occurs in the
somewhat rare minerals, _scheelite_, CaWO_{4}, and _wolfram_; the latter
being an isomorphous mixture of the normal tungstates of iron and
manganese, (MnFe)WO_{4}. Molybdenum is most frequently met with as
_molybdenite_, MoS_{2}, which presents a certain resemblance to graphite
in its physical properties and softness. It also occurs, but much more
rarely, as a yellow lead ore, PbMoO_{4}. In both these forms molybdenum
occurs in the primary rocks, in granites, gneiss, &c., and in iron and
copper ores in Saxony, Sweden, and Finland. Tungsten ores are sometimes
met with in considerable masses in the primary rocks of Bohemia and
Saxony, and also in England, America, and the Urals. The preliminary
treatment of the ore is very simple; for example, the sulphide, MoS_{2},
is roasted, and thus converted into sulphurous anhydride and molybdic
anhydride, MoO_{3}, which is then dissolved in alkalis, generally in
ammonia. The ammonium molybdate is then treated with acids, when the
sparingly soluble molybdic acid is precipitated. Wolfram is treated in a
different manner. Most frequently the finely-ground ore is repeatedly
boiled with hydrochloric and nitric acids, and the resultant solutions
(of salts of manganese and iron) poured off, until the dark brown mass of
ore disappears, whilst the tungstic acid remains, mixed with silica, as
an insoluble residue; it is treated also with ammonia, and is thus
converted into soluble ammonium tungstate, which passes into solution and
yields tungstic acid when treated with acids. This hydrate is then
ignited, and leaves tungstic anhydride. The general character of molybdic
and tungstic anhydrides is analogous to that of chromic anhydride; they
are anhydrides of a feebly acid character, which easily give polyacid
salts and colloid solutions.[8 bis]
[8 bis] The atomic composition of the tungsten and molybdenum compounds
is taken as being identical with that of the compounds of sulphur
and chromium, because (1) both these metals give two oxides in
which the amounts of oxygen per given amount of metal stand in the
ratio 2 : 3; (2) the higher oxide is of the latter kind, and, like
chromic and sulphuric anhydrides, it has an acid character; (3)
certain of the molybdates are isomorphous with the sulphates; (4)
the specific heat of tungsten is 0·0334, consequently the product
of the atomic weight and specific heat is 6·15, like that of the
other elements--it is the same with molybdenum, 96·0 × 0·0722 =
6·9; (5) tungsten forms with chlorine not only compounds WCl_{6},
WCl_{5}, and WOCl_{4}, but also WO_{2}Cl_{2}, a volatile substance
the analogue of chromyl chloride, CrO_{2}Cl_{2}, and sulphuryl
chloride, SO_{2}Cl_{2}. Molybdenum gives the chlorine compounds,
MOCl_{2}, MOCl_{3}(?), MOCl_{4} (fuses at 194°, boils at 268°;
according to Debray it contains MOCl_{5}), MoOCl_{4},
MoO_{2}Cl_{2}, and MoO_{2}(OH)Cl. The existence of tungsten
hexachloride, WCl_{6}, is an excellent proof of the fact that the
type SX_{6} appears in the analogues of sulphur as in SO_{3}; (6)
the vapour density accurately determined for the chlorine compounds
MoCl_{4}, WCl_{6}, WCl_{5}, WOCl_{4} (Roscoe) leaves no doubt as to
the molecular composition of the compounds of tungsten and
molybdenum, for the observed and calculated results entirely agree.
Tungsten is sometimes called scheele in honour of Scheele, who
discovered it in 1781 and molybdenum in 1778. Tungsten is also
known as wolfram; the former name was the name given to it by
Scheele, because he extracted it from the mineral then known as
tungsten and now called scheelite, CaWO_{4}. The researches of
Roscoe, Blomstrand and others have subsequently thrown considerable
light on the whole history of the compounds of molybdenum and
tungsten.
The ammonium salts of tungsten and molybdic acids when ignited
leave the anhydrides, which resemble each other in many respects.
_Tungsten anhydride_, WO_{3}, is a yellowish substance, which only
fuses at a strong heat, and has a sp. gr. of 6·8. It is insoluble
both in water and acid, but solutions of the alkalis, and even of
the alkali carbonates, dissolve it, especially when heated, forming
alkaline salts. _Molybdic anhydride_, MoO_{3}, is obtained by
igniting the acid (hydrate) or the ammonium salt, and forms a white
mass which fuses at a red heat, and solidifies to a yellow
crystalline mass of sp. gr. 4·4; whilst on further heating in open
vessels or in a stream of air this anhydride _sublimes_ in pearly
scales--this enables it to be obtained in a tolerably pure state.
Water dissolves it in small quantities--namely, 1 part requires 600
parts of water for its solution. The hydrates of molybdic anhydride
are _soluble also in acids_ (a hydrate, H_{2}MoO_{4}, is obtained
from the nitric acid solution of the ammonium salt), which forms
one of their distinctions from the tungstic acids. But after
ignition molybdic anhydride is insoluble in acids, like tungstic
anhydride; alkalis dissolve this anhydride, easily forming
molybdates. Potassium bitartrate dissolves the anhydride with the
aid of heat. None of the acids yet considered by us form so many
different salts with one and the same base (alkali) as molybdic and
tungstic acids. The composition of these salts, and their
properties also, vary considerably. The most important discovery in
this respect was made by Marguerite and Laurent, who showed that
the salts which contain a large proportion of tungstic acid are
easily soluble in water, and ascribed this property to the fact
that tungstic acid may be obtained _in several states_. The common
tungstates, obtained with an excess of alkali, have an alkaline
reaction, and on the addition of sulphuric or hydrochloric acid
first deposit an acid salt and then a hydrate of tungstic acid,
which is insoluble both in water and acids; but if instead of
sulphuric or hydrochloric acids, we add acetic or phosphoric acid,
or if the tungstate be saturated with a fresh quantity of tungstic
acid, which may be done by boiling the solution of the alkali salt
with the precipitated tungstic acid, a solution is obtained which,
on the addition of sulphuric or a similar acid, does not give a
precipitate of tungstic acid at the ordinary or at higher
temperatures. The solution then contains peculiar salts of tungstic
acid, and if there be an excess of acid it also contains tungstic
acid itself; Laurent, Riche, and others called it _metatungstic
acid_, and it is still known by this name. Those salts which with
acids immediately give the insoluble tungstic acid have the
composition R_{2}WO_{4},RHWO_{4}, whilst those which give the
soluble metatungstic acid contain a far greater proportion of the
acid elements. Scheibler obtained the (soluble) metatungstic acid
itself by treating the soluble barium (meta) tetratungstate,
BaO,4WO_{3}, with sulphuric acid. Subsequent research showed the
existence of a similar phenomenon for molybdic acid. There is no
doubt that this is a case of colloidal modifications.
Many chemists have worked on the various salts formed by molybdic
and tungstic acids. The tungstates have been investigated by
Marguerite, Laurent, Marignac, Riche, Scheibler, Anthon, and
others. The molybdates were partially studied by the same chemists,
but chiefly by Struvé and Svanberg, Delafontaine, and others. It
appears that for a given amount of base the salts contain one to
eight equivalents of molybdic or tungstic anhydride; _i.e._ if the
base have the composition RO, then the highest proportion of base
will be contained by the salts of the composition ROWO_{3} or
ROMoO_{3}--that is, by those salts which correspond with the normal
acids H_{2}WO_{4} and H_{2}MoO_{4}, of the same nature as sulphuric
acid; but there also exist salts of the composition RO,2WO_{3},
RO,3WO_{3} ... RO,8WO_{3}. The water contained in the composition
of many of the acid salts is often not taken into account in the
above. The properties of the salts holding different proportions of
acids vary considerably, but one salt may be converted into another
by the addition of acid or base with great facility, and the
greater the proportion of the elements of the acid in a salt, the
more stable, within a certain limit, is its solution and the salt
itself.
The most common ammonium molybdate has the composition
(NH_{4}HO)_{6},H_{2}O,7MoO_{3} (or, according to Marignac and
others, NH_{4}HMoO_{4}), and is prepared by evaporating an
ammoniacal solution of molybdic acid. It is used in the laboratory
for precipitating phosphoric acid, and is purified for this purpose
by mixing its solution with a small quantity of magnesium nitrate,
in order to precipitate any phosphoric acid present, filtering, and
then adding nitric acid and evaporating to dryness. A pure ammonium
molybdate free from phosphoric acid may then be extracted from the
residue.
Phosphoric acid forms insoluble compounds with the oxides of
uranium and iron, tin, bismuth, &c., having feeble basic and even
acid properties. This perhaps depends on the fact that the atoms of
hydrogen in phosphoric acid are of a very different character, as
we saw above. Those atoms of hydrogen which are replaced with
difficulty by ammonium, sodium, &c., are probably easily replaced
by feebly energetic acid groups--that is, the formation of
particular complex substances may be expected to take place at the
expense of these atoms of the hydrogen of phosphoric acid and of
certain feeble metallic acids; and these substances will still be
acids, because the hydrogen of the phosphoric acids and metallic
acids, which is easily replaced by metals, is not removed by their
mutual combination, but remains in the resultant compound. Such a
conclusion is verified in the _phosphomolybdic acids_ obtained
(1888) by Debray. If a solution of ammonium molybdate be acidified,
and a small amount of a solution (it may be acid) containing
orthophosphoric acid or its salts be added to it (so that there are
at least 40 parts of molybdic acid present to 1 part of phosphoric
acid), then after a period of twenty-four hours the whole of the
phosphoric acid is separated as a yellow precipitate, containing,
however, not more than 3 to 4 p.c. of phosphoric anhydride, about 3
p.c. of ammonia, about 90 p.c. of molybdic anhydride, and about 4
p.c. of water. The formation of this precipitate is so distinct and
so complete that this method is employed for the discovery and
separation of the smallest quantities of phosphoric acid.
Phosphoric acid was found to be present in the majority of rocks by
this means. The precipitate is soluble in ammonia and its salts, in
alkalis and phosphates, but is perfectly insoluble in nitric,
sulphuric, and hydrochloric acids in the presence of ammonium
molybdate. The composition of the precipitate appears to vary under
the conditions of its precipitation, but its nature became clear
when the acid corresponding with it was obtained. If the
above-described yellow precipitate be boiled in aqua regia, the
ammonia is destroyed, and an acid is obtained in solution, which,
when evaporated in the air, crystallises out in yellow oblique
prisms of approximately the composition
P_{2}O_{5},20MoO_{3},26H_{2}O. Such an unusual proportion of
component parts is explained by the above-mentioned considerations.
We saw above that molybdic acid easily gives salts
R_{2}O_n_MoO_{3}_m_H_{2}O, which we may imagine to correspond to a
hydrate MoO_{2}(HO)_{2}_n_Mo_{3}_m_H_{2}O. And suppose that such a
hydrate reacts on orthophosphoric acid, forming water and compounds
of the composition MoO_{2}(HPO_{4})_n_MoO_{3}_m_H_{2}O or
MoO_{3}(H_{2}PO_{4})_{2}_n_MoO_{3}_m_H_{2}O; this is actually the
composition of phosphomolybdic acid. Probably it contains a portion
of the hydrogen replaceable by metals of both the acids H_{3}PO_{4}
and of H_{2}MoO_{4}. The crystalline acid above is probably
H_{3}MoPO_{7},9MoO_{3},12H_{2}O. This acid is really tribasic,
because its aqueous solution precipitates salts of potassium,
ammonium, rubidium (but not lithium and sodium) _from acid
solutions_, and gives a _yellow_ precipitate of the composition
R_{3}MoPO_{7},9MoO_{3},8H_{2}O, where R = NH_{4}. Besides these,
salts of another composition may be obtained, as would be expected
from the preceding. These salts are only stable in acid solutions
(which is naturally due to their containing an excess of acid
oxides), whilst under the action of alkalis they give _colourless_
phosphomolybdates of the composition R_{3}MoPO_{3},MoO_{2},3H_{2}O.
The corresponding salts of potassium, silver, ammonium, are easily
soluble in water and crystalline.
Phosphomolybdic acid is an example of the _complex inorganic acids_
first obtained by Marignac and afterwards generalised and studied
in detail by Gibbs. We shall afterwards meet with several examples
of such acids, and we will now turn attention to the fact that they
are usually formed by weak polybasic acids (boric, silicic,
molybdic, &c.), and in certain respects resemble the cobaltic and
such similar complex compounds, with which we shall become
acquainted in the following chapter. As an example we will here
mention certain complex compounds containing molybdic and tungstic
acids, as they will illustrate the possibility of a considerable
complexity in the composition of salts. The action of ammonium
molybdate upon a dilute solution of purpureocobaltic salts (_see_
Chapter XXII.) acidulated with acetic acid gives a salt which after
drying at 100° has the composition
Co_{2}O_{3}10NH_{3}7MoO_{3}3H_{2}O. After ignition this salt leaves
a residue having the composition 2CoO_{7}MoO_{3}. An analogous
compound is also obtained for tungstic acid, having the composition
Co_{2}O_{3}10NH_{3}10WO_{3}9H_{2}O. In this case after ignition
there remains a salt of the composition CoO_{5}WO_{3} (Carnot,
1889). Professor Kournakoff, by treating a solution of potassium
and sodium molybdates, containing a certain amount of suboxide of
cobalt, with bromine obtained salts having the composition:
3K_{2}OCo_{2}O_{3}12MoO_{3}20H_{2}O (light green) and
3K_{2}OCo_{2}O_{3}10Mo_{3}10H_{2}O (dark green). Péchard (1893)
obtained salts of the four complex phosphotungstic acids by
evaporating equivalent mixtures of solutions of phosphoric acid and
metatungstic acid (_see_ further on): phosphotrimetatungstic acid
P_{2}O_{5}12WO_{3}48H_{2}O, phosphotetrametatungstic acid
P_{2}O_{5}16WO_{3}69H_{2}O, phosphopentametatungstic acid
P_{2}O_{5}20WO_{3}H_{2}O, and phosphohexametatungstic acid
P_{2}O_{5}24WO_{3}59H_{2}O. Kehrmann and Frankel described still
more complex salts, such as:
3Ag_{2}O_{4}BaOP_{2}O_{5}22WO_{3}H_{2}O,5BaO_{2}
K_{2}OP_{2}O_{3}22WO_{3}48H_{2}O.
Analogous double salts with 22WO_{3} were also obtained with KSr,
KHg, BaHg, and NH_{4}Pb. Kehrmann (1892) considers the possibility
of obtaining an unlimited number of such salts to be a general
characteristic of such compounds. Mahom and Friedheim (1892)
obtained compounds of similar complexity for molybdic and arsenic
acids.
For tungstic acid there are known: (1) Normal salts--for example,
K_{2}WO_{4}; (2) the so-called acid salts have a composition like
3K_{2}O,7WO_{3},6H_{2}O or K_{6}H_{8}(WO_{4})_{7},2H_{2}O; (3) the
tritungstates like Na_{3}O,3WO_{3},3H_{2}O =
Na_{2}H_{4}(WO_{4})_{3},H_{2}O. All these three classes of salts
are soluble in water, but are precipitated by barium chloride, and
with acids in solution give an insoluble hydrate of tungstic acid;
whilst those salts which are enumerated below do not give a
precipitate either with acids or with the salts of the heavy
metals, because they form soluble salts even with barium and lead.
They are generally called metatungstates. They all contain water
and a larger proportion of acid elements than the preceding salts;
(4) the tetratungstates, like Na_{2}O,4WO_{3},10H_{2}O and
BaO,4WO_{3},9H_{2}O for example; (5) the octatungstates--for
example, Na_{2}O,8WO_{3},24H_{2}O. Since the metatungstates lose so
much water at 100° that they leave salts whose composition
corresponds with an acid, 3H_{2}O,4WO_{3}--that is,
H_{6}W_{4}O_{15}--whilst in the meta salts only 2 hydrogens are
replaced by metals, it is assumed, although without much ground,
that these salts contain a particular soluble metatungstic acid of
the composition H_{6}W_{4}O_{15}.
As an example we will give a short description of the sodium salts.
The normal salt, Na_{2}WO_{4}, is obtained by heating a strong
solution of sodium carbonate with tungstic acid to a temperature of
80°; if the solution be filtered hot, it crystallises in rhombic
tabular crystals, having the composition Na_{2}WO_{4},2H_{2}O,
which remain unchanged in the air and are easily soluble in water.
When this salt is fused with a fresh quantity of tungstic acid, it
gives a ditungstate, which is soluble in water and separates from
its solution in crystals containing water. The same salt is
obtained by carefully adding hydrochloric acid to the solution of
the normal salt so long as a precipitate does not appear, and the
liquid still has an alkaline reaction. This salt was first supposed
to have the composition Na_{2}W_{2}O_{7},4H_{2}O, but it has since
been found to contain (at 100°) Na_{6}W_{7}O_{24},16H_{2}O--that
is, it corresponds with the similar salt of molybdic acid.
(If this salt be heated to a red heat in a stream of hydrogen, it
loses a portion of its oxygen, acquires a metallic lustre, and
turns a golden yellow colour, and, after being treated with water,
alkali, and acid, leaves golden yellow leaflets and cubes which are
very like gold. This very remarkable substance, discovered by
Wöhler, has, according to Malaguti's analysis, the composition
Na_{2}W_{3}O_{9}; that is, it, as it were, contains a double
tungstate of tungsten oxide, WO_{2}, and of sodium,
Na_{2}WO_{4},WO_{2}WO_{3}. The decomposition of the fused sodium
salt is best effected by finely-divided tin. This substance has a
sp. gr. 6·6; it conducts electricity like metals, and like them has
a metallic lustre. When brought into contact with zinc and
sulphuric acid it disengages hydrogen, and it becomes covered with
a coating of copper in a solution of copper sulphate in the
presence of zinc--that is, notwithstanding its complex composition
it presents to a certain extent the appearance and reactions of the
metals. It is not acted on by aqua regia or alkaline solutions, but
it is oxidised when ignited in air.)
The ditungstate mentioned above, deprived of water (having
undergone a modification similar to that of metaphosphoric acid),
after being treated with water, leaves an anhydrous, sparingly
soluble tetratungstate, Na_{2}WO_{4},3WO_{3}, which, when heated at
120° in a closed tube with water, passes into an easily soluble
metatungstate. It may therefore be said that the metatungstates are
hydrated compounds. On boiling a solution of the above-mentioned
salts of sodium with the yellow hydrate of tungstic acid they give
a solution of metatungstate, which is the hydrated tetratungstate.
Its crystals contain Na_{2}W_{4}O_{13},10H_{2}O. After the hydrate
of tungstic acid (obtained from the ordinary tungstates by
precipitation with an acid) has stood a long time in contact with a
solution (hot or cold) of sodium tungstate, it gives a solution
which is not precipitated by hydrochloric acid; this must be
filtered and evaporated over sulphuric acid in a desiccator (it is
decomposed by boiling). It first forms a very dense solution
(aluminium floats in it) of sp. gr. 3·0, and octahedral crystals of
_sodium metatungstate_, Na_{2}W_{4}O_{13},10H_{2}O, sp. gr. 3·85,
then separate. It effloresces and loses water, and at 100° only two
out of the ten equivalents of water remain, but the properties of
the salt remain unaltered. If the salt be deprived of water by
further heating, it becomes insoluble. At the ordinary temperature
one part of water dissolves ten parts of the metatungstate. The
other metatungstates are easily obtained from this salt. Thus a
strong and hot solution, mixed with a like solution of barium
chloride, gives on cooling crystals of barium metatungstate,
BaW_{4}O_{13},9H_{2}O. These crystals are dissolved without change
in water containing hydrochloric acid, and also in hot water, but
they are partially decomposed by cold water, with the formation of
a solution of metatungstic acid and of the normal barium salt
BaWO_{4}.
In order to explain the difference in the properties of the salts
of tungstic acid, we may add that a mixture of a solution of
tungstic acid with a solution of silicic acid does not coagulate
when heated, although the silicic acid alone would do so; this is
due to the formation of a silicotungstic acid, discovered by
Marignac, which presents a fresh example of a complex acid. A
solution of a tungstate dissolves gelatinous silica, just as it
does gelatinous tungstic acid, and when evaporated deposits a
crystalline salt of silicotungstic acid. This solution is not
precipitated either by acids (a clear analogy to the
metatungstates) or by sulphuretted hydrogen, and corresponds with a
series of salts. These salts contain one equivalent of silica and 8
equivalents of hydrogen or metals, in the same form as in salts, to
12 or 10 equivalents of tungstic anhydride; for example the
crystalline potassium salt has the composition
K_{8}W_{12}SiO_{42},14H_{2}O = 4K_{2}O,12WO_{3},SiO_{2},14H_{2}O.
Acid salts are also known in which half of the metal is replaced by
hydrogen. The complexity of the composition of such complex acids
(for example, of the phosphomolybdic acid) involuntarily leads to
the idea of polymerisation, which we were obliged to recognise for
silica, lead oxide, and other compounds. This polymerisation, it
seems to me, may be understood thus: a hydrate A (for example,
tungstic acid) is capable of combining with a hydrate B (for
example, silica or phosphoric acid, with or without the
disengagement of water), and by reason of this faculty it is
capable of polymerisation--that is, A combines with A--combines
with itself--just as aldehyde, C_{2}H_{4}O, or the cyanogen
compounds are able to combine with hydrogen, oxygen, &c., and are
liable to polymerisation. On this view the molecule of tungstic
acid is probably much more complex than we represent it; this
agrees with the easy volatility of such compounds as the
chloranhydrides, CrO_{2}Cl_{2}, MoO_{2}Cl_{2}, the analogues of the
volatile sulphuryl chloride, SO_{2}Cl_{2}, and with the
non-volatility, or difficult volatility, of chromic and molybdic
anhydrides, the analogues of the volatile sulphuric anhydride. Such
a view also finds a certain confirmation in the researches made by
Graham on the _colloidal_ state of tungstic acid, because colloidal
properties only appertain to compounds of a very complex
composition. The observations made by Graham on the colloidal state
of tungstic and molybdic acids introduced much new matter into the
history of these substances. When sodium tungstate, mixed in a
dilute solution with an equivalent quantity of dilute hydrochloric
acid, is placed in a dialyser, hydrochloric acid and sodium
chloride pass through the membrane, and a solution of tungstic acid
remains in the dialyser. Out of 100 parts of tungstic acid about 80
parts remain in the dialyser. The solution has a bitter, astringent
taste, and does not yield gelatinous tungstic acid (hydrogel)
either when heated or on the addition of acids or salts. It may
also be evaporated to dryness; it then forms a vitreous mass of the
_hydrosol_ of _tungstic acid_, which adheres strongly to the walls
of the vessel in which it has been evaporated, and is perfectly
soluble in water. It does not even lose its solubility after having
been heated to 200°, and only becomes insoluble when heated to a
red heat, when it loses about 2-1/2 p.c. of water. The dry acid,
dissolved in a small quantity of water, forms a gluey mass, just
like gum arabic, which is one of the representatives of the
hydrosols of colloidal substances. The solution, containing 5 p.c
of the anhydride, has a sp. gr. of 1·047; with 20 p.c., of 1·217;
with 50 p.c., of 1·80; and with 80 p.c., of 3·24. The presence of a
polymerised trioxide in the form of hydrate, H_{2}OW_{3}O_{9} or
H_{2}O_{4}WO_{3}, must then be recognised in the solution: this is
confirmed by Sabaneeff's cryoscopic determinations (1889). A
similar stable solution of molybdic acid is obtained by the
dialysis of a mixture of a strong solution of sodium molybdate with
hydrochloric acid (the precipitate which is formed is
re-dissolved). If MoCl_{4} be precipitated by ammonia and washed
with water, a point is reached at which perfect solution takes
place, and the molybdic acid forms a colloid solution which is
precipitated by the addition of ammonia (Muthmann). The addition of
alkali to the solutions of the hydrosols of tungstic and molybdic
acids immediately results in the re-formation of the ordinary
tungstates and molybdates. There appears to be no doubt but that
the same transformation is accomplished in the passage of the
ordinary tungstates into the metatungstates as takes place in the
passage of tungstic acid itself from an insoluble into a soluble
state; but this may be even actually proved to be the case, because
Scheibler obtained a solution of tungstic acid, before Graham, by
decomposing barium metatungstate (BaO_{4}WO_{3},9H_{2}O) with
sulphuric acid. By treating this salt with sulphuric acid in the
amount required for the precipitation of the baryta, Scheibler
obtained a solution of metatungstic acid which, when containing
43·75 p.c. of acid, had a sp. gr. of 1·634, and with 27·61 p.c. a
sp. gr. of 1·327--that is, specific gravities corresponding with
those found by Graham.
Péchard found that as much heat is evolved by neutralising
metatungstic acid as with sulphuric acid.
Questions connected with the metamorphoses or modifications of
tungstic and molybdic acids, and the polymerisation and colloidal
state of substances, as well as the formation of complex acids,
belong to that class of problems the solution of which will do much
towards attaining a true comprehension of the mechanism of a number
of chemical reactions. I think, moreover, that questions of this
kind stand in intimate connection with the theory of the formation
of solutions and alloys and other so-called indefinite compounds.
Hydrogen (which does not directly form compounds with Cr, Mo, and W)
reduces molybdic and tungstic anhydride at a red heat; and this forms the
means of obtaining metallic molybdenum and tungsten. _Both metals_ are
infusible, and both under the action of heat form compounds with carbon
and iron (the addition of tungsten to steel renders the latter ductile
and hard).[9] Molybdenum forms a grey powder, which scarcely aggregates
under a most powerful heat, and has a specific gravity of 8·5. It is not
acted on by the air at the ordinary temperature, but when ignited it is
first converted into a brown, and then into a blue oxide, and lastly into
molybdic anhydride. Acids do not act on it--that is, it does not liberate
hydrogen from them, not even from hydrochloric acid--but strong sulphuric
acid disengages sulphurous anhydride, forming a brown mass, containing a
lower oxide of molybdenum. Alkalis in solution do not act on molybdenum,
but when fused with it hydrogen is given off, which shows, as does its
whole character, the acid properties of the metal. The properties of
tungsten are almost identical; it is infusible, has an iron-grey colour,
is exceedingly hard, so that it even scratches glass. Its specific
gravity is 19·1 (according to Roscoe), so that, like uranium, platinum,
&c., it is one of the heaviest metals.[9 bis] Just as sulphur and
chromium have their corresponding persulphuric and perchromic acids,
H_{2}S_{2}O_{8} and H_{2}CrO_{8}, having the properties of peroxides, and
corresponding to peroxide of hydrogen, so also molybdenum and tungsten
are known to give _permolybdic_ and _pertungstic_ acids, H_{2}Mo_{2}O_{8}
and H_{2}W_{2}O_{8}, which have the properties of true peroxides, _i.e._
easily disengage iodine from KI and chlorine from HCl, easily part with
their oxygen, and are formed by the action of peroxide of hydrogen, into
which they are readily reconverted (hence they may be regarded as
compounds of H_{2}O_{2} with 2MoO_{3} and 2WO_{3}), &c. Their formation
(Boerwald 1884, Kemmerer 1891) is at once seen in the coloration (not
destroyed by boiling), which is obtained on mixing a solution of the
salts with peroxide of hydrogen, and on treating, for instance, molybdic
acid with a solution of peroxide of hydrogen (Péchard 1892). The acid
then forms an orange-coloured solution, which after evaporation in vacuo
leaves Mo_{2}H_{2}O_{8}4H_{2}O as a crystalline powder, and loses 4H_{2}O
at 100°, beyond which it decomposes with the evolution of oxygen.[9 tri]
[9] Moissan (1893) studied the compounds of Mo and W formed with carbon
in the electrical furnace (they are extremely hard) from a mixture
of the anhydrides and carbon. Poleck and Grützner obtained definite
compounds FeW_{2} and FeW_{2}C_{3} for tungsten. Metallic W and Mo
displace Ag from its solutions but not Pb. There is reason for
believing that the sp. gr. of pure molybdenum is higher than that
(8·5) generally ascribed to it.
[9 bis] We may conclude our description of tungsten and molybdenum by
stating that their sulphur compounds have an acid character, like
carbon bisulphide or stannic sulphide. If sulphuretted hydrogen be
passed through a solution of a molybdate it does not give a
precipitate unless sulphuric acid be present, when a dark brown
precipitate of _molybdenum trisulphide_, MoS_{3}, is formed. When
this sulphide is ignited without access of air it gives the
bisulphide MoS_{2}; the latter is not able to combine with
potassium sulphide like the trisulphide MoS_{3}, which forms a
salt, K_{2}MoS_{4}, corresponding with K_{2}MoO_{4}. This is
soluble in water, and separates out from its solution in red
crystals, which have a metallic lustre and reflect a green light.
It is easily obtained by heating the native bisulphide, MoS_{2},
with potash, sulphur, and a small amount of charcoal, which serves
for deoxidising the oxygen compounds. Tungsten gives similar
compounds, R_{2}WS_{4}, where R = NH_{4}, K, Na. They are
decomposed by acids, with the separation of tungsten trisulphide,
WS_{3}, and molybdenum trisulphide, MoS_{3}. Rideal (1892) obtained
W_{2}N_{3} by heating WO_{3} in NH_{3}. This compound exhibited the
general properties of metallic nitrides.
[9 tri] When peroxide of hydrogen acts upon a solution of potassium
molybdate well-formed yellow crystals belonging to the triclinic
system separate out in the cold. When these crystals are heated in
vacuo they first lose water and then decompose, leaving a residue
composed of the salt originally taken. They are soluble in water
but insoluble in alcohol. Their composition is represented by the
formula K_{2}Mo_{2}O_{8}2H_{2}O. An ammonium salt is obtained by
evaporating peroxide of hydrogen with ammonium molybdate. The
following salts have also been obtained by the action of peroxide
of hydrogen upon the corresponding molybdates:
Na_{2}Mo_{2}O_{6}6H_{2}O--in yellow prismatic crystals;
MgMo_{2}O_{8}10H_{2}O--stellar needles; BaMoO_{8}2H_{2}O--in
microscopic yellow octahedra. A corresponding sodium pertungstate
has been obtained by Péchard by boiling sodium tungstate with a
solution of peroxide of hydrogen for several minutes. The solution
rapidly turns yellow, and no longer gives a precipitate of tungstic
anhydride when treated with nitric acid. When evaporated in vacuo
the solution leaves a thick syrupy liquid from which ray-like
crystals separate out; these crystals are more soluble in water
than the salt originally taken. When heated they also lose water
and oxygen. Their composition answers to the formula
M_{2}W_{2}O_{8}2H_{2}O, where M = Na, NH_{4}, &c. The permolybdates
and pertungstates have similar properties. When treated with oxygen
acids they give peroxide of hydrogen, and disengage chlorine and
iodine from hydrochloric acid and potassium iodide.
Piccini (1891) showed that peroxide of hydrogen not only combines
with the oxygen compounds of Mo and W, but also with their
fluo-compounds, among which ammonium fluo-molybdate
MoO_{2}F_{2}2NH_{4} and others have long been known. (A few new
salts of similar composition have been obtained by F. Moureu in
1893.) The action of peroxide of hydrogen upon these compounds
gives salts containing a larger amount of oxygen; for instance, a
solution of MoO_{2}F_{2}2KFH_{2}O with peroxide of hydrogen gives a
yellow solution which after cooling separates out yellow
crystalline flakes of MoO_{3}F_{2}2KFH_{2}O, resembling the salt
originally taken in their external appearance. By employing a
similar method Piccini also obtained:
MoO_{3}F_{2}2RbFH_{2}O--yellow monoclinic crystals;
MoO_{3}F_{2},2CsFH_{2}O,--yellow flakes, and the corresponding
tungstic compounds. All these salts react like peroxide of
hydrogen.
In speaking of these compounds I for my part think it may be well
to call attention to the fact that, in the first place, the
composition of Piccini's oxy-fluo compounds does not correspond to
that of permolybdic and pertungstic acid. If the latter be
expressed by formulæ with one equivalent of an element, they will
be HMoO_{4} and HWO_{4}, and the oxy-fluo form corresponding to
them should have the composition MoO_{3}F and WO_{3}F while it
contains MO_{3}F_{2} and WO_{3}F_{2}, _i.e._ answers as it were to
a higher degree of oxidation, MoH_{2}O_{3} and W_{3}HO_{3}. But if
permolybdic acid be regarded as 2MoO_{3} + H_{2}O_{2}, _i.e._ as
containing the elements of peroxide of hydrogen, then Piccini's
compound will also be found to contain the original salts + H_{2}O;
for example, from MoO_{2}F_{2}2KFH_{2}O there is obtained a
compound MoO_{2}F_{2}2KFH_{2}O_{2}, _i.e._ instead of H_{2}O they
contain H_{2}O_{2}. In the second place the capacity of the salts
of molybdenum and tungsten to retain a further amount of oxygen or
H_{2}O_{2} probably bears some relation to their property of giving
complex acids and of polymerising which has been considered in Note
8 bis. There is, however, a great chemical interest in the
accumulation of data respecting these high peroxide compounds
corresponding to molybdic and tungstic acids. With regard to the
peroxide form of uranium, _see_ Chapter XX., Note 66.
_Uranium_, U = 240, has the highest atomic weight of all the analogues
of chromium, and indeed of all the elements yet known. Its highest
salt-forming oxide, UO_{3}, shows very feeble acid properties. Although
it gives sparingly-soluble yellow compounds with alkalis, which fully
correspond with the dichromates--for example, Na_{2}U_{2}O_{7} =
Na_{2}O,2UO_{3},[10]--yet it more frequently and easily reacts with
acids, HX, forming fluorescent yellowish-green salts of the composition
UO_{2}X_{2}, and in this respect uranic trioxide, UO_{3}, differs from
chromic anhydride, CrO_{3}, although the latter is able to give the
oxychloride, CrO_{2}Cl_{2}. In molybdenum and tungsten, however, we see a
clear transition from chromium to uranium. Thus, for example, chromyl
chloride, CrO_{2}Cl_{2}, is a brown liquid which volatilises without
change, and is completely decomposed by water; molybdenum oxychloride,
MoO_{2}Cl_{2}, is a crystalline substance of a yellow colour, which is
volatile and soluble in water (Blomstrand), like many salts. Tungsten
oxychloride, WO_{2}Cl_{2}, stands still nearer to uranyl chloride in its
properties; it forms yellow scales on which water and alkalis act, as
they do on many salts (zinc chloride, ferric chloride, aluminium
chloride, stannic chloride, &c.), and perfectly corresponds with the
difficultly volatile salt, UO_{2}Cl_{2} (obtained by Peligot by the
action of chlorine on ignited uranium dioxide, UO_{2}), which is also
yellow and gives a yellow solution with water, like all the salts
UO_{2}X_{2}. The property of uranic oxide, UO_{3}, of forming salts
UO_{2}X_{2} is shown in the fact that the hydrated oxide of uranium,
UO_{2}(HO)_{2}, which is obtained from the nitrate, carbonate, and other
salts by the loss of the elements of the acid, is easily soluble in
acids, as well as in the fact that the lower grades of oxidation of
uranium are able, when treated with nitric acid, to form an easily
crystallisable uranyl nitrate, UO_{2}(NO_{3})_{2},6H_{2}O; this is the
most commonly occurring uranium salt.[11]
[10] Uranium trioxide, or uranic oxide, shows its feeble basic and acid
properties in a great number of its reactions. (1) Solutions of
uranic salts give yellow precipitates with alkalis, but these
precipitates do not contain the hydrate of the oxide, but
compounds of it with bases; for example, 2UO_{2}(NO_{3})_{2} +
6KHO = 4KNO_{3} + 3H_{2}O + K_{2}U_{2}O_{7}. There are other
_urano-alkali compounds_ of the same constitution; for example,
(NH_{4})_{2}U_{2}O_{7} (known commercially as uranic oxide),
MgU_{2}O_{7}, BaU_{2}O_{7}. They are the analogues of the
dichromates. Sodium uranate is the most generally used under the
name of uranium yellow, Na_{2}U_{2}O_{7}. It is used for imparting
the characteristic yellow-green tint to glass and porcelain.
Neither heat nor water nor acids are able to extract the alkali
from sodium uranate, Na_{2}U_{2}O_{7}, and therefore it is a true
insoluble salt, of a yellow colour, and clearly indicates the acid
character (although feeble) of uranic oxide. (2) The carbonates of
the alkaline earths (for instance, barium carbonate) precipitate
uranic oxide from its salts, as they do all the salts of feeble
bases; for example, R_{2}O_{3}. (3) The _alkaline carbonates_,
when added to solutions of uranic salts, give a _precipitate,
which is soluble in_ _an excess of the reagent_, and particularly
so if the acid carbonates be taken. This is due to the fact that
(4) the uranyl salts _easily form double salts_ with the salts of
the alkali metals, including the salts of ammonium. Uranium, in
the form of these double salts, often gives salts of well-defined
crystalline form, although the simple salts are little prone to
appear in crystals. Such, for example, are the salts obtained by
dissolving potassium uranate, K_{2}U_{2}O_{7}, in acids, with the
addition of potassium salts of the same acids. Thus, with
hydrochloric acid and potassium chloride a well-formed crystalline
salt, K_{2}(UO_{2})Cl_{4},2H_{2}O, belonging to the monoclinic
system, is produced. This salt decomposes in dissolving in pure
water. Among these double salts we may mention the double
carbonate with the alkalis, R_{4}(UO_{2})(CO_{3})_{3} (equal to
2R_{2}CO_{3} + UO_{2}CO_{3}); the acetates,
R(UO_{2})(C_{2}H_{3}O_{2})_{3}--for instance, the sodium salt,
Na(UO_{2})(C_{2}H_{3}O_{2})_{3}, and the potassium salt,
K(UO_{2})(C_{2}H_{3}O_{2})_{3},H_{2}O; the sulphates,
R_{2}(UO_{2})(SO_{4})_{3},2H_{2}O, &c. In the preceding formula R
= K, Na, NH_{4}, or R_{2} = Mg, Ba, &c. _This property of giving
comparatively stable double salts indicates feebly developed basic
properties_, because double salts are mainly formed by salts of
distinctly basic metals (these form, as it were, the basic element
of a double salt) and salts of feebly energetic bases (these form
the acid element of a double salt), just as the former also give
acid salts; the acid of the acid salts is replaced in the double
salts by the salt of the feebly energetic base, which, like water,
belongs to the class of intermediate bases. For this reason barium
does not give double salts with alkalis as magnesium does, and
this is why double salts are more easily formed by potassium than
by lithium in the series of the alkali metals. (5) The most
remarkable property, proving the feeble energy of uranic oxide as
a base, is seen in the fact that when their composition is
compared with that of other salts those of uranic oxide _always
appear as basic salts_. It is well known that a normal salt,
R_{2}X_{6}, corresponds with the oxide R_{2}O_{3}, where X = Cl,
NO_{3}, &c., or X_{2} = SO_{4}, CO_{3}, &c.; but there also exist
basic salts of the same type where X = HO or X_{2} = O. We saw
salts of all kinds among the salts of aluminium, chromium, and
others. With uranic oxide no salts are known of the types UX_{6}
(UCl_{6}, U(SO_{4})_{3}, alums, &c., are not known), nor even
salts, U(HO)_{2}X_{4} or UOX_{4}, but it always forms salts of the
type U(HO)_{4}X_{2}, or UO_{2}X_{2}. Judging from the fact that
nearly all the salts of uranic oxide retain water in crystallising
from their solutions, and that this water is difficult to separate
from them, it may be thought to be water of hydration. This is
seen in part from the fact that the composition of many of the
salts of uranic oxide may then be expressed without the presence
of water of crystallisation; for instance, U(HO)_{4}K_{2}Cl_{4}
(and the salt of NH_{4}, U(HO)_{4}K_{2}(SO_{4})_{2},
U(HO)_{4}(C_{2}H_{3}O_{2})_{2}. Sodium uranyl acetate however does
not contain water.
[11] _Uranyl nitrate_, or uranium nitrate, UO_{2}(NO_{3})_{2},6H_{2}O,
crystallises from its solutions in transparent yellowish-green
prisms (from an acid solution), or in tabular crystals (from a
neutral solution), which effloresce in the air and are easily
soluble in water, alcohol, and ether, have a sp. gr. of 2·8, and
fuse when heated, losing nitric acid and water in the process. If
the salt itself (Berzelius) or its alcoholic solution (Malaguti)
be heated up to the temperature at which oxides of nitrogen are
evolved, there then remains a mass which, after being evaporated
with water, leaves uranyl hydroxide, UO_{2}(HO)_{2} (sp. gr.
5·93), whilst if the salt be ignited there remains the dioxide,
UO_{2}, as a brick-red powder, which on further heating loses
oxygen and forms the dark olive uranoso-uranic oxide, U_{3}O_{8}.
The solution of the nitrate obtained from the ore is purified in
the following manner: sulphurous anhydride is first passed through
it in order to reduce the arsenic acid present into arsenious
acid; the solution is then heated to 60°, and sulphuretted
hydrogen passed through it; this precipitates the lead, arsenic,
and tin, and certain other metals, as sulphides, insoluble in
water and dilute nitric acid. This liquid is then filtered and
evaporated with nitric acid to crystallisation, and the crystals
are dissolved in ether. Or else the solution is first treated with
chlorine in order to convert the ferrous chloride (produced by the
action of the hydrogen sulphide) into ferric chloride, the oxides
are then precipitated by ammonia, and the resultant precipitate,
containing the oxides Fe_{2}O_{3}, UO_{3}, and compounds of the
latter with potash, lime, ammonia, and other bases present in the
solution (the latter being due to the property of uranic oxide of
combining with bases), is washed and dissolved in a strong,
slightly-heated solution of ammonium carbonate, which dissolves
the uranic oxide but not the ferric oxide. The solution is
filtered, and on cooling deposits a well-crystallising _uranyl
ammonium carbonate_, UO_{2}(NH_{4})_{4}(CO_{3})_{3}, in brilliant
monoclinic crystals which on exposure to air slowly give off
water, carbonic anhydride, and ammonia; the same decomposition is
readily effected at 300°, the residue then consisting of uranic
oxide. This salt is not very soluble in water, but is readily so
in ammonium carbonate; it is obvious that it may readily be
converted into all the other salts of oxides of uranium. Uranium
salts are also purified in the form of _acetate_, which is very
sparingly soluble, and is therefore directly precipitated from a
strong solution of the nitrate by mixing it with acetic acid.
We may also mention the _uranyl phosphate_, HUPO_{6}, which must
be regarded as an orthophosphate in which two hydrogens are
replaced by the radicle uranyl, UO_{2}, _i.e._ as H(UO_{2})PO_{4}.
This salt is formed as a hydrated gelatinous yellow precipitate,
on mixing a solution of uranyl nitrate with disodium phosphate.
The precipitation occurs in the presence of acetic acid, but not
in the presence of hydrochloric acid. If moreover an excess of an
ammonium salt be present, the ammonia enters into the composition
of the bright yellow gelatinous precipitate formed, in the
proportion UO_{2}NH_{4}PO_{4}. This precipitate is not soluble in
water and acetic acid, and its solution in inorganic acids when
boiled entirely expels all the phosphoric acid. This fact is taken
advantage of for removing phosphoric acids from solutions--for
instance, from those containing salts of calcium and magnesium.
_Uranium_, which gives an oxide, UO_{3}, and the corresponding salt
UO_{2}X_{2} and dioxide UO_{2}, to which the salts UX_{4} correspond, is
rarely met with in nature. Uranite or the double orthophosphate of uranic
oxide, R(UO_{2})H_{2}P_{2}O_{8},7H_{2}O, where R = Cu or Ca,
uranium-vitriol U(SO_{4})_{2},H_{2}O, samarakite, and æschynite, are very
rarely found, and then only in small quantities. Of more frequent and
abundant occurrence is the non-crystalline, earthy brown uranium ore
known as _pitchblende_ (sp. gr. 7·2), which is mainly composed of the
intermediate oxide, U_{3}O_{8} = UO_{2},2UO_{3}. This ore is found at
Joachimsthal in Bohemia and in Cornwall. It usually contains a number of
different impurities, chiefly sulphides and arsenides of lead and iron,
as well as lime and silica compounds. In order to expel the arsenic and
sulphur it is roasted, ground, washed with dilute hydrochloric acid,
which does not dissolve the uranoso-uranic oxide, U_{3}O_{8}, and the
residue is dissolved in nitric acid, which transforms the uranium oxide
into the nitrate, UO_{2}(NO_{3})_{2}.
It must be observed that the oxide of uranium, first distinguished by
Klaproth (1789), was for a long time regarded as able to give metallic
uranium under the action of charcoal and other reducing agents (with the
aid of heat). But the substance thus obtained was only the _uranium
dioxide_, UO_{2}. The compound nature of this dioxide,[12] or the
presence of oxygen in it, was demonstrated by Peligot (1841), by igniting
it with charcoal in a stream of chlorine. He thus obtained a volatile
_uranium tetrachloride_, UCl_{4},[13] which, when heated with sodium,
gave _metallic uranium_ as a grey metal, having a specific gravity of
18·7, and liberating hydrogen from acids, with the formation of green
uranous salts, UX_{4}, which act as powerful reducing agents.[14]
[12] Uranium dioxide, or _uranyl_, UO_{2}, which is contained in the
salts UO_{2}X_{2}, has the appearance and many of the properties
of a metal. Uranic oxide may be regarded as uranyl oxide,
(UO_{2})O, its salts as salts of this uranyl; its hydroxide,
(UO_{2})H_{2}O_{2}, is constituted like CaH_{2}O_{2}. The green
oxide of uranium, uranoso-uranic oxide (easily formed from uranic
salts by the loss of oxygen), U_{3}O_{8} = UO_{2},2UO_{3}, when
ignited with charcoal or hydrogen (dry) gives a brilliant
crystalline substance of sp. gr. about 11·0 (Urlaub), whose
appearance resembles that of metals, and decomposes steam at a red
heat with the evolution of hydrogen; it does not, however,
decompose hydrochloric or sulphuric acid, but is oxidised by
nitric acid. The same substance (i.e. uranium dioxide UO_{2}) is
also obtained by igniting the compound (UO_{2})K_{2}Cl_{4} in a
stream of hydrogen, according to the equation UO_{2}K_{4}Cl_{4} +
H_{2} = UO_{2} + 2HCl + 2KCl. It was at first regarded as the
metal. In 1841 Peligot found that it contained oxygen, because
carbonic oxide and anhydride were evolved when it was ignited with
charcoal in a stream of chlorine, and from 272 parts of the
substance which was considered to be metal he obtained 382 parts
of a volatile product containing 142 parts of chlorine. From this
it was concluded that the substance taken contained an equivalent
amount of oxygen. As 142 parts of chlorine correspond with 32
parts of oxygen, it followed that 272 - 32 = 240 parts of metal
were combined in the substance with 32 parts of oxygen, and also
in the chlorine compound obtained with 142 parts of chlorine.
These calculations have been made for the now accepted atomic
weight of uranium (U = 240, _see_ Note 14). Peligot took another
atomic weight, but this does not alter the principle of the
argument.
[13] _Uranium tetrachloride_, uranous chloride, UCl_{4}, corresponds
with uranous oxide as a base. It was obtained by Peligot by
igniting uranic oxide mixed with charcoal in a stream of _dry_
chlorine: UO_{3} + 3C + 2Cl_{2} = UCl_{4} + 3CO. This green
volatile compound (Note 12) crystallises in regular octahedra, is
very hygroscopic, easily soluble in water, with the development of
a considerable amount of heat, and no longer separates out from
its solution in an anhydrous state, but disengages hydrochloric
acid when evaporated. The solution of uranous chloride in water is
green. It is also formed by the action of zinc and copper (forming
cuprous chloride) on a solution of uranyl chloride, UO_{2}Cl_{2},
especially in the presence of hydrochloric acid and sal-ammoniac.
Solutions of uranyl salts are converted into uranous salts by the
action of various reducing agents, and among others by organic
substances or by the action of light, whilst the salts UX_{4} are
converted into uranyl salts, UO_{2}X_{2}, by exposure to air or by
oxidising agents. Solutions of the green uranyl salts act as
powerful reducing agents, and give a brown precipitate of the
uranous hydroxide, UH_{4}O_{4}, with potash and other alkalis.
This hydroxide is easily soluble in acids but not in alkalis. On
ignition it does not form the oxide UO_{2}, because it decomposes
water, but when the higher oxides of uranium are ignited in a
stream of hydrogen or with charcoal they yield uranous oxide. Both
it and the chloride UCl_{4}, dissolve in strong sulphuric acid,
forming a green salt, U(SO_{4})_{2},2H_{2}O. The same salt,
together with uranyl sulphate, UO_{2}(SO_{4}), is formed when the
green oxide, U_{3}O_{8}, is dissolved in hot sulphuric acid. The
salts obtained in the latter instance may be separated by adding
alcohol to the solution, which is left exposed to the light; the
alcohol reduces the uranyl salt to uranous salt, an excess of acid
being required. An excess of water decomposes this salt, forming a
basic salt, which is also easily produced under other
circumstances, and contains UO(SO_{4}),2H_{2}O (which corresponds
to the uranic salt).
[14] The atomic weight of uranium was formerly taken as half the
present one, U = 120, and the oxides U_{2}O_{3}, suboxide UO, and
green oxide U_{3}O_{4}, were of the same types as the oxides of
iron. With a certain resemblance to the elements of the iron
group, uranium presents many points of distinction which do not
permit its being grouped with them. Thus uranium forms a very
stable oxide, U_{2}O_{3}(U = 120), but does not give the
corresponding chloride U_{2}Cl_{6} (Roscoe, however, in 1874
obtained UCl_{5}, like MoCl_{5} and WCl_{5}), and under those
circumstances (the ignition of oxide of uranium mixed with
charcoal, in a stream of chlorine), when the formation of this
compound might be expected, it gives (U = 120) the chloride
UCl_{2}, which is characterised by its volatility; this is not a
property, to such an extent, of any of the bichlorides, RCl_{2},
of the iron group.
The alteration or doubling of the atomic weight of uranium--_i.e._
the recognition of U = 240--was made for the first time in the
first (Russian) edition of this work (1871), and in my memoir of
the same year in Liebig's _Annalen_, on the ground that with an
atomic weight 120, uranium could not be placed in the periodic
system. I think it will not be superfluous to add the following
remarks on this subject: (1) In the other groups (K--Rb--Cs,
Ca--Sr--Ba, Cl--Br--I) the acid character of the oxides decreases
and their basic character increases with the rise of atomic
weight, and therefore we should expect to find the same in the
group Cr--Mo--W--U, and if CrO_{3}, MoO_{3}, WO_{3} be the
anhydrides of acids then we indeed find a decrease in their acid
character, and therefore uranium trioxide, UO_{3}, should be a
very feeble anhydride, but its basic properties should also be
very feeble. Uranic oxide does indeed show these properties, as
was pointed out above (Note 10). (2) Chromium and its analogues,
besides the oxides RO_{3}, also form lower grades of oxidation
RO_{2}, R_{2}O_{3}, and the same is seen in uranium; it forms
UO_{3}, UO_{2}, U_{2}O_{3} and their compounds. (3) Molybdenum and
tungsten, in being reduced from RO_{3}, easily and frequently give
an intermediate oxide of a blue colour, and uranium shows the same
property; giving the so-called green oxide which, according to
present views, must be regarded as U_{3}O_{8} = UO_{2}2UO_{3},
analogous to Mo_{3}O_{8}. (4) The higher chlorides, RCl_{6},
possible for the elements of this group, are either unstable
(WCl_{6}) or do not exist at all (Cr); but there is one single
lower volatile compound, which is decomposed by water, and liable
to further reduction into a non-volatile chlorine product and the
metal. The same is observed in uranium, which forms an easily
volatile chloride, UCl_{4}, decomposed by water. (5) The high sp.
gr. of uranium (18·6) is explained by its analogy to tungsten (sp.
gr. 19·1). (6) For uranium, as for chromium and tungsten, yellow
tints predominate in the form RO_{3}, whilst the lower forms are
green and blue. (7) Zimmermann (1881) determined the vapour
densities of uranous bromide, UBr_{4}, and chloride, UCl_{4} (19·4
and 13·2), and they were found to correspond to the formulæ given
above--that is, they confirmed the higher atomic weight U = 240.
Roscoe, a great authority on the metals of this group, was the
first to accept the proposed atomic weight of uranium, U = 240,
which since Zimmermann's work has been generally recognised.
As the salts of uranic oxide are reduced in the absence of organic matter
by the action of light, and as they impart a characteristic coloration to
glass,[15] they find a certain application in photography and glass work.
[15] Uranium glass, obtained by the addition of the yellow salt
K_{2}U_{2}O_{7} to glass, has a green yellow fluorescence, and is
sometimes employed for ornaments; it absorbs the violet rays, like
the other salts of uranic oxide--that is, it possesses an
absorption spectrum in which the violet rays are absent. The index
of refraction of the absorbed rays is altered, and they are given
out again as greenish-yellow rays; hence, compounds of uranic
acid, when placed in the violet portion of the spectrum, emit a
greenish-yellow light, and this forms one of the best examples
(another is found in a solution of quinine sulphate) of the
phenomenon of fluorescence. The rays of light which pass through
uranic compounds do not contain the rays which excite the
phenomena of fluorescence and of chemical transformation, as the
researches of Stokes prove.
If we compare together the highly acid elements, sulphur, selenium, and
tellurium, of the uneven series, with chromium, molybdenum, tungsten, and
uranium of the even series, we find that the resemblance of the
properties of the higher form RO_{3} does not extend to the lower forms,
and even entirely disappears in the elements, for there is not the
smallest resemblance between sulphur and chromium and their analogues in
a free state. In other words, this means that the small periods, like Na,
Mg, Al, Si, P, S, Cl, containing seven elements, do not contain any near
analogues of chromium, molybdenum, &c., and therefore their true position
among the other elements must be looked for only in those large periods
which contain two small periods, and whose type is seen in the period
containing: K, Ca, Sc, Ti, V, Cr, Mn, Fe, Co, Ni, Cu, Zn, Ga, Ge, As, Se,
Br. These large periods contain Ca and Zn, giving RO, Sc, and Ga of the
third group, Ti and Ge giving RO_{2}, V and As forming R_{2}O_{5}, Cr and
Se of the sixth group, Mn and Br of the seventh group, and the remaining
elements, Fe, Co, Ni, form connective members of the intermediate eighth
group, to the description of the representatives of which we shall turn
in the following chapters. We will now proceed to describe _manganese_,
Mn = 55, as an element of the seventh group of the even series, directly
following after Cr = 52, which corresponds with Br = 80 to the same
degree that Cr does with Se = 79. For chromium, selenium, and bromine
very close analogues are known, but for manganese as yet none have been
obtained--that is, it is the only representative of the even series in
the seventh group. In placing manganese with the halogens in one group,
the periodic system of the elements only requires that it should bear an
analogy to the halogens in the higher type of oxidation--_i.e._ in the
salts and acids--whilst it requires that as great a difference should be
expected in the lower types and elements as there exists between chromium
or molybdenum and sulphur or selenium. And this is actually the case. The
elements of the seventh group form a higher salt-forming oxide,
R_{2}O_{7}, and its corresponding hydrate, HRO_{4}, and salts--for
example, KClO_{4}. Manganese in the form of potassium permanganate,
KMnO_{4}, actually presents a great analogy in many respects to potassium
perchlorate, KClO_{4}. The analogy of the crystalline form of both salts
was shown by Mitscherlich. The salts of permanganic acid are also nearly
all soluble in water, like those of perchloric acid, and if the silver
salt of the latter, AgClO_{4}, be sparingly soluble in water, so also is
silver permanganate, AgMnO_{4}. The specific volume of potassium
perchlorate is equal to 55, because its specific gravity = 2·54; the
specific volume of potassium permanganate is equal to 58, because its
specific gravity = 2·71. So that the volumes of equivalent quantities are
in this instance approximately the same whilst the atomic volumes of
chlorine (35·5/1·3 = 27) and manganese (55/7·5) are in the ratio 4 : 1.
In a free state the higher acids HClO_{4} and HMnO_{4} are both soluble
in water and volatile, both are powerful oxidisers--in a word, their
analogy is still closer than that of chromic and sulphuric acids, and
those points of distinction which they present also appear among the
nearest analogues--for example, in sulphuric and telluric acids, in
hydrochloric and hydriodic acids, &c. Besides Mn_{2}O_{7} manganese gives
a lower grade of oxidation, MnO_{3}, analogous to sulphuric and chromic
trioxides, and with it corresponds potassium manganate, K_{2}MnO_{4},
isomorphous with potassium sulphate.[16] In the still lower grades of
oxidation, Mn_{2}O_{3} and MnO, there is hardly any similarity to
chlorine, whilst every point of resemblance disappears when we come to
the elements themselves--_i.e._ to manganese and chlorine--for manganese
is a metal, like iron, which combines directly with chlorine to form a
saline compound, MnCl_{2}, analogous to magnesium chloride.[17]
[16] The comparison of potassium permanganate with potassium
perchlorate, or of potassium manganate with potassium sulphate,
shows directly that many of the physical and chemical properties
of substances do not depend on the nature of the elements, but on
the atomic types in which they appear, on the kind of movements,
or on the positions in which the atoms forming the molecule occur.
[17] If, however, we compare the spectra (Vol. I. p. 565) of chlorine,
bromine, and iodine with that of manganese, a certain resemblance
or analogy is to be found connecting manganese both to iron and to
chlorine, bromine, and iodine.
Manganese belongs to the number of metals widely distributed in nature,
especially in those localities where iron occurs, whose ores frequently
contain compounds of manganous oxide, MnO, which presents a resemblance
to ferrous oxide, FeO, and to magnesia. In many minerals magnesia and the
oxides allied to it are replaced by manganous oxide; calcspars and
magnesites--_i.e._ R´´CO_{3} in general--are frequently met with
containing manganous carbonate, which also occurs in a separate state,
although but rarely. The soil also and the ash of plants generally
contain a small quantity of manganese. In the analysis of minerals it is
generally found that manganese occurs together with magnesia, because,
like it, manganous oxide remains in solution in the presence of
ammoniacal salts, not being precipitated by reagents. The property of
this manganous oxide, MnO, of passing into the higher grades of oxidation
under the influence of heat, alkalis, and air, gives an easy means not
only of discovering the presence of manganese in admixture with magnesia,
but also of separating these two analogous bases. Magnesia is not able to
give higher grades of oxidation, whilst manganese gives them with great
facility. Thus, for instance, an _alkaline_ solution of sodium
hypochlorite produces a precipitate of manganese dioxide in a solution of
a manganous salt: MnCl_{2} + NaClO + 2NaHO = MnO_{2} + H_{2}O + 3NaCl;
whilst magnesia is not changed under these circumstances, and remains in
the form of MgCl_{2}. If the magnesia be precipitated owing to the
presence of alkali, it may be dissolved in acetic acid, in which
manganese dioxide is insoluble. The presence of small quantities of
manganese may also be recognised by the green coloration which alkalis
acquire when heated with manganese compounds in the air. This green
coloration depends on the property of manganese of giving a green
alkaline manganate: MnCl_{2} + 4KHO + O_{2} = K_{2}MnO_{4} + 2KCl +
2H_{2}O. Thus _the faculty of oxidising in the presence of alkalis_ forms
an essential character of manganese. The higher grades of oxidation
containing Mn_{2}O_{7} and MnO_{3} are quite unknown in nature, and even
MnO_{2} is not so widely spread in nature as the ores composed of
manganous compounds which are met with nearly everywhere. The most
important ore of manganese is its dioxide, or so-called _peroxide_,
MnO_{2}, which is known in mineralogy as _pyrolusite_. Manganese also
occurs as an oxide corresponding with magnetic iron ore, MnO,Mn_{2}O_{3}
= Mn_{3}O_{4}, forming the mineral known as _hausmannite_. The oxide
Mn_{2}O_{3} also occurs in nature as the anhydrous mineral _braunite_,
and in a hydrated form, Mn_{2}O_{3},H_{2}O, called _manganite_. Both of
these often occur as an admixture in pyrolusite. Besides which, manganese
is met with in nature as a rose-coloured mineral, _rhodonite_, or
silicate, MnSiO_{3}. Very fine and rich deposits of manganese ores have
been found in the Caucasus, the Urals, and along the Dnieper. Those at
the Sharapansky district of the Government of Kutais and at Nicopol on
the Dnieper are particularly rich. A large quantity of the ore (as much
as 100,000 tons yearly) is exported from these localities.
Thus manganese gives oxides of the following forms: MnO, manganous oxide,
and manganous salts, MnX_{2}, corresponding with the base, which
resembles magnesia and ferrous oxide in many respects; Mn_{2}O_{3}, a
very feeble base, giving salts, MnX_{3}, analogous to the aluminium and
ferric salts, easily reduced to MnX_{2}; MnO_{2}, dioxide, generally
called peroxide, an almost indifferent oxide, or feebly acid;[18]
MnO_{3}, manganic anhydride, which forms salts resembling potassium
sulphate;[18 bis] Mn_{2}O_{7}, permanganic anhydride, giving salts
analogous to the perchlorates.
[18] The name 'peroxide' should only be retained for those _highest_
oxides (and MnO_{2} stands between MnO and MnO_{3}) which either
by a direct method of double decomposition are able to give
hydrogen peroxide or contain a larger proportion of oxygen than
the base or the acid, just as hydrogen peroxide contains more
oxygen than water. Their type will be H_{2}O_{2}, and they are
exemplified by barium peroxide, BaO_{2}, and sulphur peroxide,
S_{2}O_{7}, &c. Such a dioxide as MnO_{2} is, in all probability,
a salt--that is, a manganous manganate, MnO_{3}MnO, and also, as a
basic salt of a feeble base, capable of combining with alkalis and
acids. Hence the name of manganese peroxide should be abandoned,
and replaced by manganese dioxide. PbO_{2} is better termed lead
dioxide than peroxide. Bisulphide of manganese, MnS_{2},
corresponding to iron pyrites, FeS_{2}, sometimes occurs in nature
in fine octahedra (and cube combinations), for instance, in
Sicily; it is called Hauerite.
[18 bis] On comparing the manganates with the permanganates--for
example, K_{2}MnO_{4} with KMnO_{4}--we find that they differ in
composition by the abstraction of one equivalent of the metal.
Such a relation in composition produced by oxidation is of
frequent occurrence--for instance, K_{4}Fe(CN)_{6} in oxidising
gives K_{3}Fe(CN)_{6}; H_{2}SO_{4} in oxidising gives persulphuric
acid, HSO_{4}, or H_{2}S_{7}O_{8}; H_{2}O forms HO or
H_{2}O_{2}, &c.
_All the oxides of manganese when heated with acids give salts_, MnX_{2},
corresponding with the lower grade of oxidation, _manganous oxide_, MnO.
Manganic oxide, Mn_{2}O_{3}, is a feebly energetic base; it is true that
it dissolves in hydrochloric acid and gives a dark solution containing
the salt MnCl_{3}, but the latter when heated evolves chlorine and gives
a salt corresponding with manganous oxide MnCl_{2}--_i.e._ at first:
Mn_{2}O_{3} + 6HCl = 3H_{2}O + Mn_{2}Cl_{6}, and then the Mn_{2}Cl_{6}
decomposes into 2MnCl_{2} + Cl_{2}. None of the remaining higher grades
of oxidation have a basic character, but _act as oxidising agents in the
presence of acids_, disengaging oxygen and passing into salts of the
lower grade of oxidation of manganese, MnO. Owing to this circumstance,
_the manganous salts_ are often obtained; they are, for instance, left in
the residue when the dioxide is used for the preparation of oxygen and
chlorine.[19]
[19] In the preparation of oxygen from the dioxide by means of
H_{2}SO_{4}, MnSO_{4} is formed; in the preparation of chlorine
from HCl and MnO_{2}, MnCl_{2} is obtained. These two manganous
salts may be taken as examples of compounds MnX_{2}. Manganous
sulphate generally contains various impurities, and also a large
amount of iron salt (from the native MnO_{2}), from which it
cannot be freed by crystallisation. Their removal may, however, be
effected by mixing a portion of the liquid with a solution of
sodium carbonate; a precipitate of manganous carbonate is then
formed. This precipitate is collected and washed, and then added
to the remaining mass of the impure solution of manganous
sulphate; on heating the solution with this precipitate, the whole
of the iron is precipitated as oxide. This is due to the fact that
in the solution of the manganese dioxide in sulphuric acid the
whole of the iron is converted into the ferric state (because the
dioxide acts as an oxidising agent), which, as an exceedingly
feeble base precipitated by calcium carbonate and other kindred
salts, is also precipitated by manganous carbonate. After being
treated in this manner, the solution of manganous sulphate is
further purified by crystallisation. If it be a bright red colour,
it is due to the presence of higher grades of oxidation of
manganese; they may be destroyed by boiling the solution, when the
oxygen from the oxides of manganese is evolved and a very faintly
coloured solution of manganous sulphate is obtained. This salt is
remarkable for the facility with which it gives various
combinations with water. By evaporating the almost colourless
solution of _manganous sulphate_ at very low temperatures, and by
cooling the saturated solution at about 0°, crystals are obtained
containing 7 atoms of water of crystallisation, MnSO_{4},7H_{2}O,
which are isomorphous with cobaltous and ferrous sulphates. These
crystals, even at 10°, lose 5 p.c. of water, and completely
effloresce at 15°, losing about 20 p.c. of water. By evaporating a
solution of the salt at the ordinary temperature, but not above
20°, crystals are obtained containing 5 mol. H_{2}O, which are
isomorphous with copper sulphate; whilst if the crystallisation be
carried on between 20° and 30°, large transparent prismatic
crystals are formed containing 4 mol. H_{2}O (see Nickel). A
boiling solution also deposits these crystals together with
crystals containing 3 mol. H_{2}O, whilst the first salt, when
fused and boiled with alcohol, gives crystals containing 2 mol.
H_{2}O. Graham obtained a monohydrated salt by drying the salt at
about 200°. The last atom of water is eliminated with difficulty,
as is the case with all salts like MnSO_{4}nH_{2}O. The crystals
containing a considerable amount of water are rose-coloured, and
the anhydrous crystals are colourless. The solubility of
MnSO_{4},4H_{2}O (Chapter I., Note 24) per 100 parts of water is:
at 10°, 127 parts; at 37°·5, 149 parts; at 75°, 145 parts; and at
101°, 92 parts. Whence it is seen that at the boiling-point this
salt is less soluble than at lower temperatures, and therefore a
solution saturated at the ordinary temperature becomes turbid when
boiled. Manganous sulphate, being analogous to magnesium sulphate,
is decomposed, like the latter, when ignited, but it does not then
leave manganous oxide, but the intermediate oxide, Mn_{3}O_{4}. It
gives double salts with the alkali sulphates. With aluminium
sulphate it forms fine radiated crystals, whose composition
resembles that of the alums--namely,
MnAl_{2}(SO_{4})_{4},24H_{2}O. This salt is easily soluble in
water, and occurs in nature.
_Manganous chloride_, MCl_{2}, crystallises with 4 mol. H_{2}O,
like the ferrous salt, and not with 6 mol. H_{2}O like many
kindred salts--for example, those of cobalt, calcium, and
magnesium; 100 parts of water dissolve 38 parts of the anhydrous
salt at 10° and 55 parts at 62°. Alcohol also dissolves manganous
chloride, and the alcoholic solution burns with a red flame. This
salt, like magnesium chloride, readily forms double salts. A
solution of borax gives a dirty rose-coloured precipitate having
the composition MnH_{4}(BO_{3})_{2}H_{2}O, which is used as a
drier in paint-making. Potassium cyanide produces a yellowish-grey
precipitate, MnC_{2}N_{2}, with manganous salts, soluble in an
excess of the reagent, a double salt, K_{4}MnC_{6}N_{6},
corresponding with potassium ferrocyanide, being formed. On
evaporation of this solution, a portion of the manganese is
oxidised and precipitated, whilst a salt corresponding to Gmelin's
red salt, K_{3},MnC_{6}N_{6} (_see_ Chapter XXII.), remains in
solution. Sulphuretted hydrogen does not precipitate salts of
manganese, not even the acetate, but ammonium sulphide gives a
flesh-coloured precipitate, MnS; at 320° this sulphide of
manganese passes into a green variety (Antony). Oxalic acid in
strong solutions of manganous salts gives a white precipitate of
the oxalate, MnC_{2}O_{4}. This precipitate is insoluble in water,
and is used for the preparation of manganous oxide itself because
it decomposes like oxalic acid when ignited (in a tube without
access of air), with the formation of carbonic anhydride, carbonic
oxide, and manganous oxide. _Manganous oxide_ thus obtained is a
green powder, which however oxidises with such facility that it
burns in air when brought into contact with an incandescent
substance, and passes into the red intermediate oxide Mn_{3}O_{4}.
In solutions of manganous salts, alkalis produce a precipitate of
the hydroxide MnH_{2}O_{2}, which rapidly absorbs oxygen in the
presence of air and gives the brown intermediate oxide, or, more
correctly speaking, its hydrate.
Manganous oxide, besides being obtained by the above-described
method from manganous oxalate, may also be obtained by igniting
the higher oxides in a stream of hydrogen, and also from manganese
carbonate. The manganous oxide ignited in the presence of hydrogen
acquires a great density, and is no longer so easily oxidised. It
may also be obtained in a crystalline form, if during the ignition
of the carbonate or higher oxide a trace of dry hydrochloric acid
gas be passed into the current of hydrogen. It is thus obtained in
the form of transparent emerald green crystals of the regular
system, and in this state is easily soluble in acids.
Manganous oxide in oxidising gives the _red oxide of manganese_,
Mn_{5}O_{4}. This is the most stable of all the oxides of
manganese; it is not only stable at the ordinary but also at a
high temperature--that is, it does not absorb or disengage oxygen
spontaneously. When ignited, all the higher oxides of manganese
pass into it by losing oxygen, and manganous oxide by absorbing
oxygen. This oxide does not give any distinct salts, but it
dissolves in sulphuric acid, forming a dark red solution, which
contains both manganous and manganic (of the _oxide_, Mn_{2}O_{3})
sulphates. The latter with potassium sulphate gives a manganese
alum, in which the alumina is replaced by its isomorphous oxide of
manganese. But this alum, like the solution of the intermediate
oxide in sulphuric acid, evolves oxygen and leaves a manganous
salt when slightly heated.
_Manganese dioxide_ is still less basic than the oxide, and
disengages oxygen or a halogen in the presence of acids, forming
manganous salts, like the oxide. However, if it be suspended in
ether, and hydrochloric acid gas passed into the mixture, which is
kept cool, the ether acquires a green colour, owing to the
formation of tetrachloride of manganese, MnCl_{4}, corresponding
with the dioxide which passes into solution. It is however very
unstable, being exceedingly easily decomposed with the evolution
of chlorine. The corresponding fluoride, MnF_{4}, obtained by
Nicklés is much more stable. At all events, manganese dioxide does
not exhibit any well-defined basic character, but has rather an
acid character, which is particularly shown in the compounds
MnF_{4} and MnCl_{4} just mentioned, and in the property of
manganese dioxide of combining with alkalis. If the higher grades
of oxidation of manganese be deoxidised in the presence of
alkalis, they frequently give the dioxide combined with the
alkali--for example, in the presence of potash a compound is
formed which contains K_{2}O,5MnO_{2}, which shows the weak acid
character of this oxide. When ignited in the presence of sodium
compounds manganese dioxide frequently forms Na_{2}O,8MnO_{2} and
Na_{2}O,12MnO_{2}, and lime when heated with MnO_{2} gives from
CaO,3MnO_{2} to (CaO)_{2},MnO_{2} (Rousseau) according to the
temperature. Besides which, perhaps, MnO_{2} is a saline compound,
containing MnOMnO_{3} or (MnO)_{3}Mn_{2}O_{7}, and there are
reactions which support such a view (Spring, Richards, Traube, and
others); for instance it is known that manganous chloride and
potassium permanganate give the dioxide in the presence of
alkalis.
Manganese dioxide may be obtained from manganous salts by the
action of oxidising agents. If manganous hydroxide or carbonate be
shaken up in water through which chlorine is passed, the
hypochlorite of the metal is not formed, as is the case with
certain other oxides, but manganese dioxide is precipitated:
2MnO_{2}H_{2} + Cl_{2} = MnCl_{2} + MnO_{2},H_{2}O + H_{2}O. Owing
to this fact, hypochlorites in the presence of alkalis and acetic
acid when added to a solution of manganous salts give hydrated
manganese dioxide, as was mentioned above. Manganous nitrate also
leaves manganese dioxide when heated to 200°. It is also obtained
from manganous and manganic salts of the alkalis, when they are
decomposed in the presence of a small amount of acid; the
practical method of converting the salts MnX_{2} into the higher
grades of oxidation is given in Chapter II., Note 6.
As the salts of manganous oxide MnX_{2} closely resemble (and are
isomorphous with) the salts of magnesia MgX_{2} in many respects (with
the exception of the fact that MnX_{2} are rose coloured and are easily
oxidised in the presence of alkalis), we will not dwell upon them, but
limit ourselves to illustrating the chemical character of manganese by
describing the metal and its corresponding acids. The fact alone that the
oxides of manganese are not reduced to the metal when ignited in hydrogen
(whilst the oxides of iron give metallic iron under these circumstances),
but only to manganous oxide, MnO, shows that manganese has a considerable
affinity for oxygen--that is, it is difficult to reduce. This may be
effected, however, by means of charcoal or sodium at a very high
temperature. A mixture of one of the oxides of manganese with charcoal or
organic matter gives fused _metallic manganese_ under the powerful heat
developed by coke with an artificial draught. The metal was obtained for
the first time in this manner by Gahn, after Pott, and more especially
Scheele, had in the last century shown the difference between the
compounds of iron and manganese (they were previously regarded as being
the same). Manganese is prepared by mixing one of its oxides in a
finely-divided state with oil and soot; the resultant mass is then first
ignited in order to decompose the organic matter, and afterwards strongly
heated in a charcoal crucible. The manganese thus obtained, however,
contains, as a rule, a considerable amount of silicon and other
impurities. Its specific gravity varies between 7·2 and 8·0. It has a
light grey colour, a feebly metallic lustre, and although it is very hard
it can be scratched by a file. It rapidly oxidises in air, being
converted into a black oxide; water acts on it with the evolution of
hydrogen--this decomposition proceeds very rapidly with boiling water,
and if the metal contain carbon.[20]
[20] Other chemists have obtained manganese by different methods, and
attributed different properties to it. This difference probably
depends on the presence of carbon in different proportions.
Deville obtained manganese by subjecting the pure dioxide, mixed
with pure charcoal (from burnt sugar), to a strong heat in a lime
crucible until the resultant metal fused. The metal obtained had a
rose tint, like bismuth, and like it was very brittle, although
exceedingly hard. It decomposed water at the ordinary temperature.
Brunner obtained manganese having a specific gravity of about 7·2,
which decomposed water very feebly at the ordinary temperature,
did not oxidise in air, and was capable of taking a bright polish,
like steel; it had the grey colour of cast iron, was very brittle,
and hard enough to scratch steel and glass, like a diamond.
Brunner's method was as follows: He decomposed the manganese
fluoride (obtained as a soluble compound by the action of
hydrofluoric acid on manganese carbonate) with sodium, by mixing
these substances together in a crucible and covering the mixture
with a layer of salt and fluor spar; after which the crucible was
first gradually heated until the reaction began, and then strongly
heated in order to fuse the metal separated. Glatzel (1889)
obtained 25 grms. of manganese, having a grey colour and sp. gr.
7·39, by heating a mixture of 100 grms. of MnCl_{2} with 200 grms.
KCl and 15 grms. Mg to a bright white heat. Moissan and others, by
heating the oxides of manganese with carbon in the electric
furnace, obtained carbides of manganese--for example, Mn_{3}C--and
remarked that the metal volatilised in the heat of the voltaic
arc. Metallic manganese is, however, not prepared on a large
scale, but only its alloys with carbon (they readily and rapidly
oxidise) and _ferro-manganese_ or a coarsely crystalline alloy of
iron, manganese and carbon, which is smelted in blast-furnaces
like pig-iron (_see_ Chapter XXII.) This ferro-manganese is
employed in the manufacture of steel by Bessemer's and other
processes (see Chapter XXII.) and for the manufacture of manganese
bronze. However, in America, Green and Wahl (1895) obtained almost
pure metallic manganese on a large scale. They first treat the ore
of MnO_{2} with 30 p.c. sulphuric acid (which extracts all the
oxides of iron present in the ore), and then heat it in a reducing
flame to convert it into MnO, which they mix with a powder of Al,
lime and CaF_{2} (as a flux), and heat the mixture in a crucible
lined with magnesia; a reaction immediately takes place at a
certain temperature, and a metal of specific gravity 7·3 is
obtained, which only contains a small trace of iron.
Manganese gives two compounds with _nitrogen_, Mn_{5}N_{2} and
Mn_{3}N_{2}. They were obtained by Prelinger (1894) from the
amalgam of manganese Mn_{2}Hg_{5} (obtained on a mercury anode by
the action of an electric current upon a solution of MnCl_{2});
the mercury may be removed from this amalgam by heating it in an
atmosphere of hydrogen, and then metallic manganese is obtained as
a grey porous mass of specific gravity 7·42. If this amalgam be
heated in dry nitrogen it gives Mn_{5}N_{2} (grey powder, sp. gr.
6·58), but if heated in an atmosphere of NH_{3} it gives (as also
does Mn_{5}N_{2}) Mn_{3}N_{2}, (a dark mass with a metallic
lustre, sp. gr. 6·21), which, when heated in nitrogen is converted
into Mn_{5}N_{2}, and if heated in hydrogen evolves NH_{3} and
disengages hydrogen from a solution of NH_{4}Cl. At all events,
manganese is a metal which decomposes water more easily than iron,
nickel, and cobalt.
It has been shown above that if manganese dioxide, or any lower oxide of
manganese, be heated with an alkali in the presence of air, the mixture
absorbs oxygen,[21] and forms an alkaline manganate of a green colour:
2KHO + MnO_{2} + O = K_{2}MnO_{4} + H_{2}O. Steam is disengaged during
the ignition of the mixture, and if this does not take place there is no
absorption of oxygen. The oxidation proceeds much more rapidly if, before
igniting in air, potassium chlorate or nitre be added to the mixture, and
this is the method of preparing _potassium manganate_, K_{2}MnO_{4}. The
resultant mass dissolved in a small quantity of water gives a dark green
solution, which, when evaporated under the receiver of an air-pump over
sulphuric acid, deposits green crystals of exactly the same form as
potassium sulphate--namely, six-sided prisms and pyramids. The
composition of the product is not changed by being redissolved, if
perfectly pure water free from air and carbonic acid be taken. But in the
presence of even very feeble acids the solution of this salt changes its
colour and becomes red, and deposits manganese dioxide. The same
decomposition takes place when the salt is heated with water, but when
diluted with a large quantity of unboiled water manganese dioxide does
not separate, although the solution turns red. This change of colour
depends on the fact that potassium manganate, K_{2}MnO_{4}, whose
solution is green, is transformed into potassium permanganate, KMnO_{4},
whose solution is of a red colour. The reaction proceeding under the
influence of acids and a large quantity of water is expressed in the
following manner: 3K_{2}MnO_{4} + 2H_{2}O = 2KMnO_{4} + MnO_{2} + 4KHO.
If there is a large proportion of acid and the decomposition is aided by
heat, the manganese dioxide and potassium permanganate are also
decomposed, with formation of manganous salt. Exactly the same
decomposition as takes place under the action of acids is also
accomplished by magnesium sulphate, which reacts in many cases like an
acid. When water holding atmospheric oxygen in solution acts on a
solution of potassium manganate, the oxygen combines directly with the
manganate and forms potassium permanganate, without precipitating
manganese dioxide, 2K_{2}MnO_{4} + O + H_{2}O = 2KMnO_{4} + 2KHO. Thus a
solution of potassium manganate undergoes a very characteristic change in
colour and passes from green to red; hence this salt received the name of
_chameleon mineral_.[22]
[21] Volume I. p. 157, Note 7.
[22] It was known to the alchemists by this name, but the true
explanation of the change in colour is due to the researches of
Chevillot, Edwards, Mitscherlich, and Forchhammer. The change in
colour of potassium manganate is due to its instability and to its
splitting up into two other manganese compounds, a higher and a
lower: 3MnO_{3} = Mn_{2}O_{7} + MnO_{2}. Manganese trioxide is
really decomposed in this manner by the action of water (see
later): 3MnO_{3} + H_{2}O = 2MnHO_{4} + MnO_{2} (Franke, Thorpe,
and Humbly). The instability of the salt is proved by the fact of
its being deoxidised by organic matter, with the formation of
manganese dioxide and alkali, so that, for instance, a solution of
this salt cannot be filtered through paper. The presence of an
excess of alkali increases the stability of the salt; when heated
it breaks up in the presence of water, with the evolution of
oxygen.
The method of preparing _potassium permanganate_ will be
understood from the above. There are many recipes for preparing
this substance, as it is now used in considerable quantities both
for technical and laboratory purposes. But in all cases the
essence of the methods is one and the same: a mixture of alkali
with any oxide of manganese (even manganous hydroxide, which may
be obtained from manganous chloride) is first heated in the
presence of air or of an oxidising substance (for the sake of
rapidity, with potassium chlorate); the resultant mass is then
treated with water and heated, when manganese dioxide is
precipitated and potassium permanganate remains in solution. This
solution may be boiled, as the liquid will contain free alkali;
but the solution cannot be evaporated to dryness, because a strong
solution, as well as the solid salt, is decomposed by heat.
By adding a dilute solution of manganous sulphate to a boiling
mixture of lead dioxide and dilute nitric acid, the whole of the
manganese may be converted into permanganic acid (Crum).
_Potassium permanganate_, KMnO_{4}, crystallises in well-formed, long
red prisms with a bright green metallic lustre. In the arts the potash is
frequently replaced by soda, and by other alkaline bases, but no salt of
permanganic acid crystallises so well as the potassium salt, and
therefore this salt is exclusively used in chemical laboratories. One
part of the crystalline salt dissolves in 15 parts of water at the
ordinary temperature. The solution is of a very deep _red colour_, which
is so intense that it is still clearly observable after being highly
diluted with water. In a solid state it is decomposed by heat, with
evolution of oxygen, a residue consisting of the lower oxides of
manganese and potassium oxide being left.[22 bis] A mixture of
permanganate of potassium, phosphorous and sulphur takes fire when struck
or rubbed, a mixture of the permanganate with carbon only takes fire when
heated, not when struck. The instability of the salt is also seen in the
fact that its solution is decomposed by peroxide of hydrogen, which at
the same time it decomposes. A number of substances reduce potassium
permanganate to manganese dioxide (in which case the red solution becomes
colourless).[23] Many organic substances (although far from all, even
when boiled in a solution of permanganate) act in this manner, being
oxidised at the expense of a portion of its oxygen. Thus, a solution of
sugar decomposes a cold solution of potassium permanganate. In the
presence of an excess of alkali, with a small quantity of sugar, the
reduction leads to the formation of potassium manganate, because
2KMnO_{4} + 2KHO = O + 2K_{2}MnO_{4} + H_{2}O. With a considerable amount
of sugar and a more prolonged action, the solution turns brown and
precipitates manganese dioxide or even oxide. In the oxidation of many
organic bodies by an alkaline solution of KMnO_{4} generally
three-eighths of the oxygen in the salt are utilised for oxidation:
2KMnO_{4} = K_{2}O + 2MnO_{2} + O_{3}. A portion of the alkali liberated
is retained by the manganese dioxide, and the other portion generally
combines with the substance oxidised, because the latter most frequently
gives an acid with an excess of alkali. A solution of potassium iodide
acts in a similar manner, being converted into potassium iodate at the
expense of the three atoms of oxygen disengaged by two molecules of
potassium permanganate.
[22 bis] The solution of this salt with an excess of impure commercial
alkali generally acquires a green tint.
[23] A solution of potassium permanganate gives a beautiful absorption
spectrum (Chapter XIII.) If the light in passing through this
solution loses a portion of its rays in it (if one may so account
for it), this is partially explained by the increased oxidising
power which the solution then acquires. We may here also remark
that a dilute solution of permanganate of potassium forms a
colourless solution with nickel salts, because the green colour of
the solution of nickel salts is complementary to the red. Such a
decolorised solution, containing a large proportion of nickel and
a small proportion of manganese, decomposes after a time, throws
down a precipitate, and re-acquires the green colour proper to the
nickel salts. The addition of a solution of a cobalt salt
(rose-red) to the nickel salt also destroys the colour of both
salts.
_In the presence of acids, potassium permanganate acts as an oxidising
agent_ with still greater energy than in the presence of alkalis. At any
rate, a greater proportion of oxygen is then available for oxidation,
namely, not 3/8, as in the presence of alkalis, but 5/8, because in the
first instance manganese dioxide is formed, and in the second case
manganous oxide, or rather the salt, MnX_{2}, corresponding with it.
Thus, for instance, in the presence of an excess of sulphuric acid, the
decomposition is accomplished in the following manner: 2KMnO_{4} +
3H_{2}SO_{4} = K_{2}SO_{4} + 2MnSO_{4} + 3H_{2}O + 5O. This
decomposition, however, does not proceed directly on mixing a solution of
the salt with sulphuric acid, and crystals of the salt even dissolve in
oil of vitriol without the evolution of oxygen, and this solution only
decomposes by degrees after a certain time. This is due to the fact that
sulphuric acid liberates free permanganic acid from the permanganate,[24]
which acid is stable in solution. But if, in the presence of acids and a
permanganate, there is a substance capable of absorbing oxygen--for
instance, capable of passing into a higher grade of oxidation--then the
reduction of the permanganic acid into manganous oxides sometimes
proceeds directly at the ordinary temperature. This reduction is very
clearly seen, because the solutions of potassium permanganate are red
whilst the manganous salts are almost colourless. Thus, for instance,
nitrous acid and its salts are converted into nitric acid and decolorise
the acid solution of the permanganate. Sulphurous anhydride and its salts
immediately decolorise potassium permanganate, forming sulphuric acid.
Ferrous salts, and in general salts of lower grades of oxidation capable
of being oxidised in solution, act in exactly the same manner.
Sulphuretted hydrogen is also oxidised to sulphuric acid; even mercury is
oxidised at the expense of permanganic acid, and decolorises its
solution, being converted into mercuric oxide. Moreover, the end point of
these reactions may easily be seen, and therefore, having first
determined the amount of active oxygen in one volume of a solution of
potassium permanganate, and knowing how many volumes are required to
effect a given oxidation, it is easy to determine the amount of an
oxidisable substance in a solution from the amount of permanganate
expended (Marguerite's method).
[24] If sulphuric acid is allowed to act on potassium permanganate
without any special precautions, a large amount of oxygen is
evolved (it may even explode and inflame), and a violet spray of
the decomposing permanganic acid is given off. But if the pure
salt (_i.e._ free from chlorine) be dissolved in pure well-cooled
sulphuric acid, without any rise in temperature, a green-coloured
liquid settles at the bottom of the vessel. This liquid does not
contain any sulphuric acid, and consists of permanganic anhydride,
Mn_{2}O_{7} (Aschoff, Terreil). It is impossible to prepare any
considerable quantity of the anhydride by this method, as it
decomposes with an explosion as it collects, evolving oxygen and
leaving red oxide of manganese. _Permanganic anhydride_,
Mn_{2}O_{7}, in dissolving in sulphuric acid, gives a green
solution, which (according to Franke, 1887) contains a compound
Mn_{2}SO_{10} = (MnO_{3})_{2}SO_{4}--that is, sulphuric acid in
which both hydrogens are replaced by the group MnO_{3}, which is
combined with OK in permanganate of potassium. This mixture with a
small quantity of water gives Mn_{2}O_{7}, according to the
equation: (MnO_{3})_{2}SO_{4} + H_{2}O = H_{2}SO_{4} +
Mn_{2}O_{7}, and when heated to 30° it gives _manganese trioxide_,
(MnO_{3})_{2}SO_{4} + H_{2}O = 2MnO_{2} + H_{2}SO_{4} + O. Pure
manganese trioxide is obtained if the solution of
(MnO_{3})_{2}SO_{4} be poured in drops on to sodium carbonate.
Then, together with carbonic anhydride, a spray of manganese
trioxide passes over, which may be collected in a well-cooled
receiver, and this shows that the reaction proceeds according to
the equation: (MnO_{3})_{2}SO_{4} + Na_{2}CO_{3} = Na_{2}SO_{4} +
2MnO_{3} + CO_{2} + O (Thorpe). The trioxide is decomposed by
water, forming manganese dioxide and a solution of _permanganic
acid_: 3MnO_{3} + H_{2}O = MnO_{2} + 2HMnO_{4}. The same acid is
obtained by dissolving permanganic anhydride in water.