temperature, but in the majority of cases between 0° and 100°; only
in the case of bromine the specific heat is taken (for the solid
state) at a temperature below -7°, according to Regnault's
determination. _The variation of the specific heat with a change of
temperature_ is a very complex phenomenon, the consideration of
which I think would here be out of place. I will only cite a few
figures as an example. According to Bystrom, the specific heat of
iron at 0° = 0·1116, at 100° = 0·1114, at 200° = 0·1188, at 300° =
0·1267, and at 1,400° = 0·4031. Between these last limits of
temperature a change takes place in iron (a spontaneous heating,
_recalescence_), as we shall see in Chapter XXII. For quartz
SiO_{2} Pionchon gives Q = 0·1737 + 394_t_10^{-6}-27_t_^{2}10^{-9}
up to 400°, for metallic aluminium (Richards, 1892) at 0° 0·222, at
20° 0·224, at 100° 0·232; consequently, as a rule, the specific
heat varies slightly with the temperature. Still more remarkable
are H. E. Weber's observations on the great variation of the
specific heat of charcoal, the diamond and boron:
0° 100° 200° 600° 900°
Wood charcoal 0·15 0·23 0·29 0·44 0·46
Diamond 0·10 0·19 0·22 0·44 0·45
Boron 0·22 0·29 0·35 -- --
These determinations, which have been verified by Dewar, Le
Chatelier (Chapter VIII., Note 13), Moissan, and Gauthier, the
latter finding for boron AQ = 6 at 400°, are of especial importance
as confirming the universality of Dulong and Petit's law, because
the elements mentioned above form exceptions to the general rule
when the mean specific heat is taken for temperatures between 0°
and 100°. Thus in the case of the diamond the product of A × Q at
0° = 1·2, and for boron = 2·4. But if we take the specific heat
towards which there is evidently a tendency with a rise of
temperature, we obtain a product approaching to 6 as with other
elements. Thus with the diamond and charcoal, it is evident that
the specific heat tends towards 0·47, which multiplied by 12 gives
5·6, the same as for magnesium and aluminium. I may here direct the
reader's attention to the fact that for solid elements having a
small atomic weight, the specific heat varies considerably if we
take the average figures for temperatures 0° to 100°:
Li = 7 Be = 9 B = 11 C = 12
Q = 0·94 0·42 0·24 0·20
AQ = 6·6 3·8 2·6 2·4
It is therefore clear that the specific heat of beryllium
determined at a low temperature cannot serve for establishing its
atomicity. On the other hand, the low atomic heat of charcoal,
graphite, and the diamond, boron, &c., may perhaps depend on the
complexity of the molecules of these elements. The necessity for
acknowledging a great complexity of the molecules of carbon was
explained in Chapter VIII. In the case of sulphur the molecule
contains at least S_{6} and its atomic heat = 32 × 0·163 = 5·22,
which is distinctly below the normal. If a large number of atoms of
carbon are contained in the molecule of charcoal, this would to a
certain extent account for its comparatively small atomic heat.
With respect to the specific heat of compounds, it will not be out
of place to mention here the conclusion arrived at by Kopp, that
the molecular heat (that is, the product of MQ) may be looked on as
the sum of the atomic heats of its component elements; but as this
rule is not a general one, and can only be applied to give an
approximate estimate of the specific heats of substances, I do not
think it necessary to go into the details of the conclusions
described in Liebig's 'Annalen Supplement-Band,' 1864, which
includes a number of determinations made by Kopp.
It is seen from this that the product of the specific heat of the element
into the atomic weight is an almost constant quantity, which is nearly 6.
Hence it is possible to determine the valency by the specific heats of
the metals. Thus, for instance, the specific heats of lithium, sodium,
and potassium convince us of the fact that their atomic weights are
indeed those which we chose, because by multiplying the specific heats
found by experiment by the corresponding atomic weights we obtain the
following figures: Li, 6·59, Na, 6·75 and K, 6·47. Of the alkaline earth
metals the specific heats have been determined: of magnesium = 0·245
(Regnault and Kopp), of calcium = 0·170 (Bunsen), and of barium = 0·05
(Mendeléeff). If the same composition be ascribed to the compounds of
magnesium as to the corresponding compounds of potassium, then the
equivalent of magnesium will be equal to 12. On multiplying this atomic
weight by the specific heat of magnesium, we obtain a figure 2·94, which
is half that which is given by the other solid elements and therefore the
atomic weight of magnesium must be taken as equal to 24 and not to 12.
Then the atomic heat of magnesium = 24 × 0·245 = 5·9; for calcium, giving
its compounds a composition CaX_{2}--for example CaCl_{2}, CaSO_{4}, CaO
(Ca = 40)--we obtain an atomic heat = 40 × 0·17 = 6·8, and for barium it
is equal to 137 × 0·05 = 6·8; that is, they must be counted as bivalent,
or that their atom replaces H_{2}, Na_{2}, or K_{2}. This conclusion may
be confirmed by a method of analogy, as we shall afterwards see. The
application of the principle of specific heats to the determination of
the magnitudes of the atomic weights of those metals, the magnitude of
whose atomic weights could not be determined by Avogadro-Gerhardt's law,
was made about 1860 by the Italian professor Cannizzaro.
Exactly the same conclusions respecting the bivalence of magnesium and
its analogues are obtained by comparing the specific heats of their
compounds, especially of the halogen compounds as the most simple,
with the specific heats of the corresponding alkali compounds. Thus,
for instance, the specific heats of magnesium and calcium chlorides,
MgCl_{2} and CaCl_{2}, are 0·194 and 0·164, and of sodium and potassium
chlorides, NaCl and KCl, 0·214 and 0·172, and therefore their molecular
heats (or the products QM, where M is the weight of the molecule)
are 18·4 and 18·2, 12·5 and 12·8, and hence the atomic heats (or the
quotient of QM by the number of atoms) are all nearly 6, as with the
elements. Whilst if, instead of the actual atomic weights Mg = 24 and
Ca = 40, their equivalents 12 and 20 be taken, then the atomic heats of
the chlorides of magnesium and calcium would be about 4·6, whilst those
of potassium and sodium chlorides are about 6·3.[5] We must remark,
however, that as the specific heat or the amount of heat required to
raise the temperature of a unit of weight one degree[6] is a complex
quantity--including not only the increase of the energy of a substance
with its rise in temperature, but also the external work of expansion[7]
and the internal work accomplished in the molecules causing them to
decompose according to the rise of temperature[8]--therefore it is
impossible to expect in the magnitude of the specific heat the great
simplicity of relation to composition which we see, for instance, in
the density of gaseous substances. Hence, although the specific heat
is one of the important means for determining the atomicity of the
elements, still the mainstay for a true judgment of atomicity is only
given by Avogadro-Gerhardt's law, _i.e._ this other method can only be
accessory or preliminary, and when possible recourse should be had to
the determination of the vapour density.
[5] It must be remarked that in the case of oxygen (and also hydrogen
and carbon) compounds the quotient of MQ/_n_, where _n_ is the
number of atoms in the molecule, is always less than 6 for solids;
for example, for MgO = 5·0, CuO = 5·1, MnO_{2} = 4·6, ice (Q =
0·504) = 3, SiO_{2} = 3·5, &c. At present it is impossible to say
whether this depends on the smaller specific heat of the atom of
oxygen in its solid compounds (Kopp, Note 4) or on some other
cause; but, nevertheless, taking into account this decrease
depending on the presence of oxygen, a reflection of the atomicity
of the elements may to a certain extent be seen in the specific
heat of the oxides. Thus for alumina, Al_{2}O_{3} (Q = 0·217), MQ =
22·3, and therefore the quotient MQ/_n_ = 4·5, which is nearly that
given by magnesium oxide, MgO. But if we ascribe the same
composition to alumina, as to magnesia--that is, if aluminium were
counted as divalent--we should obtain the figure 3·7, which is much
less. In general, in compounds of identical atomic composition and
of analogous chemical properties the molecular heats MQ are nearly
equal, as many investigators have long remarked. For example, ZnS =
11·7 and HgS = 11·8; MgSO_{4} = 27·0 and ZnSO_{4} = 28·0, &c.
[6] If W be the amount of heat contained in a mass _m_ of a substance
at a temperature _t_, and _d_W the amount expended in heating it
from _t_ to _t_ + _dt_, then the specific heat Q = _d_W(_m_ ×
_dt_). The specific heat not only varies with the composition and
complexity of the molecules of a substance, but also with the
temperature, pressure, and physical state of a substance. Even for
gases the variation of Q with _t_ is to be observed. Thus it is
seen from the experiments of Regnault and Wiedemann that the
specific heat of carbonic anhydride at 0° = 0·19, at 100° = 0·22,
and at 200° = 0·24. But the variation of the specific heat of
permanent gases with the temperature is, as far as we know, very
inconsiderable. According to Mallard and Le Chatelier it is =
0·0006/M per 1°, where M is the molecular weight (for instance, for
O_{2}, M = 32). Therefore the specific heat of those permanent
gases which contain two atoms in the molecule (H_{2}, O_{2}, N_{2},
CO, and NO) may be, as is shown by experiment, taken as not varying
with the temperature. The constancy of the specific heat of perfect
gases forms one of the fundamental propositions of the whole theory
of heat and on it depends the determination of temperatures by
means of gas-thermometers containing hydrogen, nitrogen, or air. Le
Chatelier (1887), on the basis of existing determinations,
concludes that the molecular heat--that is, the product MQ--of all
gases varies in proportion to the temperature, and tends to become
equal (= 6·8) at the temperature of absolute zero (that is,
at -273°); and therefore MQ = 6·8 + _a_(273 + _t_), where _a_ is a
constant quantity which increases with the complexity of the
gaseous molecule and Q is the specific heat of the gas under a
constant pressure. For permanent gases _a_ almost = 0, and
therefore MQ = 6·8--that is, the atomic heat (if the molecule
contains two atoms) = 3·4, as it is in fact (Chapter IX., Note
17 bis). As regards liquids (as well as the vapours formed by
them), the specific heat always rises with the temperature. Thus
for benzene it equals 0·38 + 0·0014_t_. R. Schiff (1887) showed
that the variation of the specific heat of many organic liquids is
proportional to the change of temperature (as in the case of gases,
according to Le Chatelier), and reduced these variations into
dependence with their composition and absolute boiling point. It is
very probable that the theory of liquids will make use of these
simple relations which recall the simplicity of the variation of
the specific gravity (Chapter II., Note 34), cohesion, and other
properties of liquids with the temperature. They are all expressed
by the linear function of the temperature, _a_ + _bt_, with the
same degree of proximity as the property of gases is expressed by
the equation _pv_ = _Rt_.
As regards the relation between the specific heats of liquids (or
of solids) and of their vapours, the specific heat of the vapour
(and also of the solid) is always less than that of the liquid. For
example, benzene vapour 0·22, liquid 0·38; chloroform vapour 0·13,
liquid 0·23; steam 0·475, liquid water 1·0. But the complexity of
the relations existing in specific heat is seen from the fact that
the specific heat of ice = 0·502 is less than that of liquid water.
According to Regnault, in the case of bromine the specific heat of
the vapour = 0·055 at (150°), of the liquid = 0·107 (at 30°), and
of solid bromine = 0·084 (at -15°). The specific heat of solid
benzoic acid (according to experiment and calculation, Hess, 1888)
between 0° and 100° is 0·31, and of liquid benzoic acid 0·50. One
of the problems of the present day is the explanation of those
complex relations which exist between the composition and such
properties as specific heat, latent heat, expansion by heat,
compression, internal friction, cohesion, and so forth. They can
only be connected by a complete theory of liquids, which may now
soon be expected, more especially as many sides of the subject have
already been partially explained.
[7] According to the above reasons the quantity of heat, Q, required to
raise the temperature of one part by weight of a substance by one
degree may be expressed by the sum Q = K + B + D, where K is the
heat actually expended in heating the substance, or what is termed
the absolute specific heat, B the amount of heat expended in the
internal work accomplished with the rise of temperature, and D the
amount of heat expended in external work. In the case of gases the
last quantity may be easily determined, knowing their coefficient
of expansion, which is approximately = 0·00368. By applying to this
case the same argument given at the end of Note 11, Chapter I., we
find that one cubic metre of a gas heated 1° produces an external
work of 10333 × 0·00368, or 38·02 kilogrammetres, on which
38·02/424 or 0·0897 heat units are expended. This is the heat
expended for the external work produced by one cubic metre of a
gas, but the specific heat refers to units of weight, and therefore
it is necessary in order to know D to reduce the above quantity to
a unit of weight. One cubic metre of hydrogen at 0° and 760 mm.
pressure weighs 0·0896 kilo, a gas of molecular weight M has a
density M/2, consequently a cubic metre weighs (at 0° and 760 mm.)
0·0448M kilo, and therefore 1 kilogram of the gas occupies a volume
1/0·0448M cubic metres, and hence the external work D in the
heating of 1 kilo of the given gas through 1° = 0·0896/0·0448M, or
D = 2/M.
Taking the magnitude of the internal work B for gases as negligible
if permanent gases are taken, and therefore supposing B = 0, we
find the specific heat of gases at a constant pressure Q = K + 2 M,
where K is the specific heat at a constant volume, or the true
specific heat, and M the molecular weight. Hence K = Q-2/M. The
magnitude of the specific heat Q is given by direct experiment.
According to Regnault's experiments, for oxygen it = 0·2175, for
hydrogen 3·405, for nitrogen 0·2438; the molecular weights of these
gases are 32, 2, and 28, and therefore for oxygen K = 0·2175-0·0625
= 0·1550, for hydrogen K = 3·4050-1·000 = 2·4050, and for nitrogen
K = 0·2438-0·0714 = 0·1724. These true specific heats of elements
are in inverse proportion to their atomic weights--that is, their
product by the atomic weight is a constant quantity. In fact, for
oxygen this product = 0·155 × 16 = 2·48, for hydrogen 2·40, for
nitrogen 0·7724 × 14 = 2·414, and therefore if A stand for the
atomic weight we obtain the expression K × A = a constant, which
may be taken as 2·45. This is the true expression of Dulong and
Petit's law, because K is the true specific heat and A the weight
of the atom. It should be remarked, moreover, that the product of
the observed specific heat Q into A is also a constant quantity
(for oxygen = 3·48, for hydrogen = 3·40), because the external work
D is also inversely proportional to the atomic weight.
In the case of gases we distinguish the specific heat at a constant
pressure _c´_ (we designated this quantity above by Q), and at a
constant volume _c_. It is evident that _the relation between the
two specific heats, k_, judging from the above, is the ratio of Q
to K, or equal to the ratio of 2·45_n_ + 2 to 2·45_n_. When _n_ = 1
this ratio _k_ = 1·8; when _n_ = 2, _k_ = 1·4, when _n_ = 3, _k_ =
1·3, and with an exceedingly large number _n_, of atoms in the
molecule, _k_ = 1. That is, the ratio between the specific heats
decreases from 1·8 to 1·0 as the number of atoms, _n_, contained in
the molecule increases. This deduction is verified to a certain
extent by direct experiment. For such gases as hydrogen, oxygen,
nitrogen, carbonic oxide, air, and others in which _n_ = 2, the
magnitude of _k_ is determined by methods described in works on
physics (for example, by the change of temperature with an
alteration of pressure, by the velocity of sound, &c.) and is found
in reality to be nearly 1·4, and for such gases as carbonic
anhydride, nitric dioxide, and others it is nearly 1·3. Kundt and
Warburg (1875), by means of the approximate method mentioned in
Note 29, Chapter VII., determined _k_ for mercury vapour when _n_ =
1, and found it to be = 1·67--that is, a larger quantity than for
air, as would be expected from the above.
It may be admitted that the true atomic heat of gases = 2·43, only
under the condition that they are distant from a liquid state, and
do not undergo a chemical change when heated--that is, when no
internal work is produced in them (B = 0). Therefore this work may
to a certain extent be judged by the observed specific heat. Thus,
for instance, for chlorine (Q = 0·12, Regnault; _k_ = 1·33,
according to Straker and Martin, and therefore K = 0·09, MK = 6·4),
the atomic heat (3·2) is much greater than for other gases
containing two atoms in a molecule, and it must be assumed,
therefore, that when it is heated some great internal work is
accomplished.
In order to generalise the facts concerning the specific heat of
gases and solids, it appears to me possible to accept the following
general proposition: _the atomic heat_ (that is, AQ or QM/_n_,
where M is the molecular weight and _n_ the number of molecules) is
_smaller_ (in solids it attains its highest value 6·8 and in gases
3·4), _the more complex the molecule_ (i.e. _the greater the number
(n) of atoms forming it_) _and so much smaller, up to a certain
point_ (in similar physical states) _the smaller the mean atomic
weight M/n_.
[8] As an example, it will be sufficient to refer to the specific heat
of nitrogen tetroxide, N_{2}O_{4}, which, when heated, gradually
passes into NO_{2}--that is, chemical work of decomposition
proceeds, which consumes heat. Speaking generally, specific heat is
a complex quantity, in which it is clear that thermal data (for
instance, the heat of reaction) alone cannot give an idea either of
chemical or of physical changes individually, but always depend on
an association of the one and the other. If a substance be heated
from _t__{0} to _t__{1} it cannot but suffer a chemical change
(that is, the state of the atoms in the molecules changes more or
less in one way or another) if dissociation sets in at a
temperature _t__{1}. Even in the case of the elements whose
molecules contain only one atom, a true chemical change is possible
with a rise of temperature, because more heat is evolved in
chemical reactions than that quantity which participates in purely
physical changes. One gram of hydrogen (specific heat = 3·4 at a
constant pressure) cooled to the temperature of absolute zero will
evolve altogether about one thousand units of heat, 8 grams of
oxygen half this amount, whilst in combining together they evolve
in the formation of 9 grams of water more than thirty times as much
heat. Hence the store of chemical energy (that is, of the motion of
the atoms, vortex, or other) is much greater than the physical
store proper to the molecules, but it is the change accomplished by
the former that is the cause of chemical transformations. Here we
evidently touch on those limits of existing knowledge beyond which
the teaching of science does not yet allow us to pass. Many new
scientific discoveries have still to be made before this is
possible.
Among the bivalent metals the first place, with respect to their
distribution in nature, is occupied by _magnesium_ and _calcium_, just as
sodium and potassium stand first amongst the univalent metals. The
relation which exists between the atomic weights of these four metals
confirms the above comparison. In fact, the combining weight of magnesium
is equal to 24, and of calcium 40; whilst the combining weights of sodium
and potassium are 23 and 39--that is, the latter are one unit less than
the former.[9] They all belong to the number of _light metals_, as they
have but a small specific gravity, in which respect they differ from the
ordinary, generally known heavy, or ore, metals (for instance, iron,
copper, silver, and lead), which are distinguished by a much greater
specific gravity. There is no doubt that their low specific gravity has a
significance, not only as a simple point of distinction, but also as a
property which determines the fundamental properties of these metals.
Indeed, all the light metals have a series of points of resemblance with
the metals of the alkalis; thus both magnesium and calcium, like the
metals of the alkalis, decompose water (without the addition of acids),
although not so easily as the latter metals. The process of the
decomposition is essentially one and the same; for example, Ca + 2H_{2}O
= CaH_{2}O_{2} + H_{2}--that is, hydrogen is liberated and a hydroxide of
the metal formed. These hydroxides are bases which neutralise nearly all
acids. However, the hydroxides RH_{2}O_{2} of calcium and magnesium are
in no respect so energetic as the hydroxides of the true metals of the
alkalis; thus when heated they lose water, are not so soluble, develop
less heat with acids, and form various salts, which are less stable and
more easily decomposed by heat than the corresponding salts of sodium and
potassium. Thus calcium and magnesium carbonates easily part with
carbonic anhydride when ignited; the nitrates are also very easily
decomposed by heat, calcium and magnesium oxides, CaO and MgO, being left
behind. The chlorides of magnesium and calcium, when heated with water,
evolve hydrogen chloride, forming the corresponding hydroxides, and when
ignited the oxides themselves. All these points are evidence of a
weakening of the alkaline properties.
[9] As if NaH = Mg and KH = Ca, which is in accordance with their
valency. KH includes two monovalent elements, and is a bivalent
group like Ca.
These metals have been termed _the metals of the alkaline earths_,
because they, like the alkali metals, form energetic bases. They are
called alkaline _earths_ because they are met with in nature in a state
of combination, forming the insoluble mass of the earth, and because as
oxides, RO, they themselves have an earthy appearance. Not a few salts of
these metals are known which are insoluble in water, whilst the
corresponding salts of the alkali metals are generally soluble--for
example, the carbonates, phosphates, borates, and other salts of the
alkaline earth metals are nearly insoluble. This enables us to separate
the metals of the alkaline earths from the metals of the alkalis. For
this purpose a solution of ammonium carbonate is added to a mixed
solution of salts of both kinds of metals, when by a double decomposition
the insoluble carbonates of the metals of the alkaline earths are formed
and fall as a precipitate, whilst the metals of the alkalis remain in
solution: RX_{2} + Na_{2}CO_{3} = RCO_{3} + 2NaX.
We may here remark that the oxides of the metals of the alkaline earths
are frequently called by special names: MgO is called magnesia or bitter
earth; CaO, lime; SrO, strontia; and BaO, baryta.
In the primary rocks the oxides of calcium and magnesium are combined
with silica, sometimes in variable quantities, so that in some cases
the lime predominates and in other cases the magnesium. The two oxides,
being analogous to each other, replace each other in equivalent
quantities. The various forms of _augite_, _hornblende_, or _amphibole_,
and of similar minerals, which enter into the composition of nearly
all rocks, contain lime and magnesia and silica. The majority of the
primary rocks also contain alumina, potash, and soda. These rocks,
under the action of water (containing carbonic acid) and air, give
up lime and magnesia to the water, and therefore they are contained
in all kinds of water, and especially in sea-water. The _carbonates_
CaCO_{3} and MgCO_{3}, frequently met with in nature, _are soluble
in an excess of water saturated with carbonic anhydride_,[10] and
therefore many natural waters contain these salts, and are able to
yield them when evaporated. However, one kilogram of water saturated
with carbonic anhydride does not dissolve more than three grams of
calcium carbonate. By gradually expelling the carbonic anhydride from
such water, an insoluble precipitate of calcium carbonate separates
out. It may confidently be stated that the formation of the very widely
distributed strata of calcium and magnesium carbonates was of this
nature, because these strata are of a sedimentary character--that is,
such as would be exhibited by a gradually accumulating deposit on the
bottom of the sea, and, moreover, frequently containing the remains of
marine plants, and animals, shells, &c. It is very probable that the
presence of these organisms in the sea has played the chief part in the
precipitation of the carbonates from the sea water, because the plants
absorb CO_{2}, and many of the organisms CaCO_{3}, and after death give
deposits of carbonate of lime; for instance, chalk, which is almost
entirely composed of the minute remains of the calcareous shields of
such organisms. These deposits of calcium and magnesium carbonates are
the most important sources of these metals. Lime generally predominates,
because it is present in rocks and running water in greater quantity
than magnesia, and in this case these sedimentary rocks are termed
_limestone_. Some common flagstones used for paving, &c., and chalk may
be taken as examples of this kind of formation. Those limestones in
which a considerable portion of the calcium is replaced by magnesium are
termed _dolomites_. The dolomites are distinguished by their hardness,
and by their not parting with the whole of their carbonic anhydride
so easily as the limestones under the action of acids. Dolomites[11]
sometimes contain an equal number of molecules of calcium carbonate and
magnesium carbonate, and they also sometimes appear in a crystalline
form, which is easily intelligible, because calcium carbonate itself
is exceedingly common in this form in nature, and is then known as
_calc spar_, whilst natural crystalline magnesium carbonate is termed
_magnesite_. The formation of the crystalline varieties of the insoluble
carbonates is explained by the possibility of a slow deposition from
solutions containing carbonic acid. Besides which (Chapter X.) calcium
and magnesium sulphates are obtained from sea water, and therefore they
are met with both as deposits and in springs. It must be observed that
magnesium is held in considerable quantities in sea water, because the
sulphate and chloride of magnesium are very soluble in water, whilst
calcium sulphate is but little soluble, and is used in the formation
of shells; and therefore if the occurrence of considerable deposits of
magnesium sulphate cannot be expected in nature, still, on the other
hand, one would expect (and they do actually occur) large masses of
calcium sulphate or _gypsum_, CaSO_{4},2H_{2}O. Gypsum sometimes forms
strata of immense size, which extend over many hectometres--for example,
in Russia on the Volga, and in the Donetz and Baltic provinces.
[10] Sodium carbonate and other carbonates of the alkalis give acid
salts which are less soluble than the normal; here, on the
contrary, with an excess of carbonic anhydride, a salt is formed
which is more soluble than the normal, but this acid salt is more
unstable than sodium hydrogen carbonate, NaHCO_{3}.
[11] The formation of dolomite may be explained, if only we imagine
that a solution of a magnesium salt acts on calcium carbonate.
Magnesium carbonate may be formed by double decomposition, and it
must be supposed that this process ceases at a certain limit
(Chapter XII.), when we shall obtain a mixture of the carbonates
of calcium and magnesium. Haitinger heated a mixture of calcium
carbonate, CaCO_{3}, with a solution of an equivalent quantity of
magnesium sulphate, MgSO_{4}, in a closed tube at 200°, and then a
portion of the magnesia actually passed into the state of
magnesium carbonate, MgCO_{3}, and a portion of the lime was
converted into gypsum, CaSO_{4}. Lubavin (1892) showed that
MgCO_{3} is more soluble than CaCO_{3} in salt water, which is of
some significance in explaining the composition of sea water.
Lime and magnesia also, but in much smaller quantities (only to the
amount of several fractions of a per cent. and rarely more), enter into
the composition of every fertile soil, and without these bases the soil
is unable to support vegetation. Lime is particularly important in this
respect, and its presence in a larger quantity generally improves the
harvest, although purely calcareous soils are as a rule infertile. For
this reason the soil is fertilised both with lime[12] itself and with
marl--that is, with clay mixed with a certain quantity of calcium
carbonate, strata of which are found nearly everywhere.
[12] The undoubted action of lime in increasing the fertility of
soils--if not in every case, at all events, with ordinary soils
which have long been under corn--is based not so much on the need
of plants for the lime itself as on those chemical and physical
changes which it produces in the soil, as a particularly powerful
base which aids the alteration of the mineral and organic elements
of the soil.
From the soil the lime and magnesia (in a smaller quantity) pass into the
substance of _plants_, where they occur as salts. Certain of these salts
separate in the interior of plants in a crystalline form--for example,
calcium oxalate. The lime occurring in plants serves as the source for
the formation of the various calcareous secretions which are so common in
_animals_ of all classes. The bones of the highest animal orders, the
shells of mollusca, the covering of the sea-urchin, and similar solid
secretions of sea animals, contain calcium salts; namely, the shells
mainly calcium carbonate, and the bones mainly calcium phosphate. Certain
limestones are almost entirely formed of such deposits. Odessa is
situated on a limestone of this kind, composed of shells. Thus magnesium
and calcium occur throughout the entire realm of nature, but calcium
predominates.
As lime and magnesia form bases which are in many respects analogous,
they were not distinguished from each other for a long time. Magnesia
was obtained for the first time in the seventeenth century from Italy,
and used as a medicine; and it was only in the last century that Black,
Bergmann, and others distinguished magnesia from lime.
_Metallic magnesium_ (and calcium also) is not obtained by heating
magnesium oxide or the carbonate with charcoal, as the alkali metals
are obtained,[13] but is liberated by the action of a galvanic current
on fused magnesium chloride (best mixed with potassium chloride); Davy
and Bussy obtained metallic magnesium by acting on magnesium chloride
with the vapours of potassium. At the present time (Deville's process)
magnesium is prepared in rather considerable quantities by a similar
process, only the potassium is replaced by sodium. Anhydrous magnesium
chloride, together with sodium chloride and calcium fluoride, is fused
in a close crucible. The latter substances only serve to facilitate the
formation of a fusible mass before and after the reaction, which is
indispensable in order to prevent the access and action of air. One part
of finely divided sodium to five parts of magnesium chloride is thrown
into the strongly heated molten mass, and after stirring the reaction
proceeds very quickly, and magnesium separates, MgCl_{2} + Na_{2} = Mg
+ 2NaCl. In working on a large scale, the powdery metallic magnesium
is then subjected to distillation at a white heat. The distillation
of the magnesium is necessary, because the undistilled metal is not
homogeneous[14] and burns unevenly: the metal is prepared for the
purpose of illumination. Magnesium is a white metal, like silver; it
is not soft like the alkali metals, but is, on the contrary, hard like
the majority of the ordinary metals. This follows from the fact that it
melts at a somewhat high temperature--namely, about 500°--and boils at
about 1000°. It is malleable and ductile, like the generality of metals,
so that it can be drawn into wires and rolled into ribbon; it is most
frequently used for lighting purposes in the latter form. Unlike the
alkali metals, magnesium does not decompose the atmospheric moisture at
the ordinary temperature, so that it is almost unacted on by air; it is
not even acted on by water at the ordinary temperature, so that it may
be washed to free it from sodium chloride. Magnesium only decomposes
water with the evolution of hydrogen at the boiling point of water,[15]
and more rapidly at still higher temperatures. This is explained by the
fact that in decomposing water magnesium forms an insoluble hydroxide,
MgH_{2}O_{2}, which covers the metal and hinders the further action
of the water. Magnesium easily displaces hydrogen from acids, forming
magnesium salts. When ignited it _burns_, not only in oxygen but in air
(and even in carbonic anhydride), forming a white powder of magnesium
oxide, or magnesia; in burning it emits a white and exceedingly
_brilliant light_. The strength of this light naturally depends on the
fact that magnesium (24 parts by weight) in burning evolves about
140 thousand heat units, and that the product of combustion, MgO,
is infusible by heat; so that the vapour of the burning magnesium
contains an ignited powder of non-volatile and infusible magnesia,
and consequently presents all the conditions for the production of
a brilliant light. The light emitted by burning magnesium contains
many rays which act chemically, and are situated in the violet and
ultra-violet parts of the spectrum. For this reason burning magnesium
may be employed for producing photographic images.[16]
[13] Sodium and potassium only decompose magnesium oxide at a white
heat and very feebly, probably for two reasons. In the first
place, because the reaction Mg + O develops more heat (about 140
thousand calories) than K_{2} + O or Na_{2} + O (about 100
thousand calories); and, in the second place, because magnesia is
not fusible at the heat of a furnace and cannot act on the
charcoal, sodium, or potassium--that is, it does not pass into
that mobile state which is necessary for reaction. The first
reason alone is not sufficient to explain the absence of the
reaction between charcoal and magnesia, because iron and charcoal
in combining with oxygen evolve less heat than sodium or
potassium, yet, nevertheless, they can displace them. With respect
to magnesium chloride, it acts on sodium and potassium, not only
because their combination with chlorine evolves more heat than the
combination of chlorine and magnesium (Mg + Cl_{2} gives 150 and
Na_{2} + Cl_{2} about 195 thousand calories), but also because a
fusion, both of the magnesium chloride and of the double salt,
takes place under the action of heat. It is probable, however,
that a reverse reaction will take place. A reverse reaction might
probably be expected, and Winkler (1890) showed that Mg reduces
the oxides of the alkali metals (Chapter XIII., Note 42).
[14] Commercial magnesium generally contains a certain amount of
magnesium nitride (Deville and Caron), Mg_{3}N_{2}--that is, a
product of substitution of ammonia which is directly formed (as is
easily shown by experiment) when magnesium is heated in nitrogen.
It is a yellowish green powder, which gives ammonia and magnesia
with water, and cyanogen when heated with carbonic anhydride.
Pashkoffsky (1893) showed that Mg_{3}N_{2} is easily formed and is
the sole product when Mg is heated to redness in a current of
NH_{3}. Perfectly pure magnesium may be obtained by the action of
a galvanic current.
[15] Hydrogen peroxide (Weltzien) dissolves magnesium. The reaction has
not been investigated.
[16] A special form of apparatus is used for burning magnesium. It is a
clockwork arrangement in which a cylinder rotates, round which a
ribbon or wire of magnesium is wound. The wire is subjected to a
uniform unwinding and burning as the cylinder rotates, and in this
manner the combustion may continue uniform for a certain time. The
same is attained in special lamps, by causing a mixture of sand
and finely divided magnesium to fall from a funnel-shaped
reservoir on to the flame. In photography it is best to blow
finely divided magnesium into a colourless (spirit or gas) flame,
and for instantaneous photography to light a cartridge of a
mixture of magnesium and chlorate of potassium by means of a spark
from a Ruhmkorff's coil (D. Mendeléeff, 1889).
Owing to its great affinity for oxygen, magnesium _reduces_ many metals
(zinc, iron, bismuth, antimony, cadmium, tin, lead, copper, silver, and
others) from solutions of their salts at the ordinary temperature,[17]
and at a red heat finely divided magnesium takes up the oxygen from
silica, alumina, boric anhydride, &c.; so that silicon and similar
elements may be obtained by directly heating a mixture of powdered silica
and magnesium in an infusible glass tube.[18]
[17] According to the observations of Maack, Comaille, Böttger, and
others. The reduction by heat mentioned further on was pointed out
by Geuther, Phipson, Parkinson and Gattermann.
[18] This action of metallic magnesium in all probability depends,
although only partially (_see_ Note 13), on its volatility, and on
the fact that, in combining with a given quantity of oxygen, it
evolves more heat than aluminium, silicon, potassium, and other
elements.
The affinity of magnesium for the halogens is much more feeble than for
oxygen,[19] as is at once evident from the fact that a solution of iodine
acts feebly on magnesium; still magnesium burns in the vapours of iodine,
bromine, and chlorine. The character of magnesium is also seen in the
fact that all its salts, especially in the presence of water, are
decomposable at a comparatively moderate temperature, the elements of the
acid being evolved, and the magnesium oxide, which is non-volatile and
unchangeable by heat, being left. This naturally refers to those acids
which are themselves volatilised by heat. Even magnesium sulphate is
completely decomposed at the temperature at which iron melts, oxide of
magnesium remaining behind. This decomposition of magnesium salts by heat
proceeds much more easily than that of calcium salts. For example,
magnesium carbonate is totally decomposed at 170°, magnesium oxide being
left behind. This _magnesia_, or _magnesium oxide_, is met with both in
an anhydrous and hydrated state in nature (the anhydrous magnesia as the
mineral _periclase_, MgO, and the hydrated magnesia as _brucite_,
MgH_{2}O_{2}). Magnesia is a well-known medicine (calcined
magnesia--_magnesia usta_). It is a white, extremely fine, and very
voluminous powder, of specific gravity 3·4; it is infusible by heat, and
only shrinks or shrivels in an oxyhydrogen flame. After long contact the
anhydrous magnesia combines with water, although very slowly, forming the
hydroxide Mg(HO)_{2}, which, however, parts with its water with great
ease when heated even below a red heat, and again yields anhydrous
magnesia. This hydroxide is obtained directly as a gelatinous amorphous
substance when a soluble alkali is mixed with a solution of any magnesium
salt, MgCl_{2} + 2KHO = Mg(HO)_{2} + 2KCl. This decomposition is
complete, and nearly all the magnesium passes into the precipitate; and
this clearly shows the almost perfect insolubility of magnesia in water.
Water dissolves a scarcely perceptible quantity of magnesium
hydroxide--namely, one part is dissolved by 55,000 parts of water. Such a
solution, however, has an alkaline reaction, and gives, with a salt of
phosphoric acid, a precipitate of magnesium phosphate, which is still
more insoluble. Magnesia is not only dissolved by acids, forming salts,
but it also displaces certain other bases--for example, ammonia from
ammonium salts when boiled; and the hydroxide also absorbs carbonic
anhydride from the air. The magnesium salts, like those of calcium,
potassium, and sodium, are colourless if they are formed from colourless
acids. Those which are soluble have a bitter taste, whence magnesia has
been termed _bitter-earth_. In comparison with the alkalis magnesia is a
feeble base, inasmuch as it forms somewhat unstable salts, easily gives
basic salts, forms acid salts with difficulty, and is able to give double
salts with the salts of the alkalis, which facts are characteristic of
feeble bases, as we shall see in becoming acquainted with the different
metals.
[19] Davy, on heating magnesia in chlorine, concluded that there was a
complete substitution, because the volume of the oxygen was half
the volume of the chlorine; it is probable, however, that owing to
the formation of chlorine oxide (Chapter XI., Note 30) the
decomposition is not complete and is limited by a reverse
reaction.
The power of magnesium salts to form double and basic salts is very
frequently shown in reactions, and is specially marked as
regards ammonium salts. If saturated solutions of magnesium and
ammonium sulphates are mixed together, a crystalline double salt
Mg(NH_{4})_{2}(SO_{4})_{2},6H_{2}O,[20] is immediately precipitated.
A strong solution of ordinary ammonium carbonate dissolves magnesium
oxide or carbonate, and precipitates crystals of a double salt,
Mg(NH_{4})_{2}(CO_{3})_{2},4H_{2}O, from which water extracts the
ammonium carbonate. With an excess of an ammonium salt the double
salt passes into solution,[21] and therefore if a solution contain
a magnesium salt and an excess of an ammonium salt--for instance,
sal-ammoniac--then sodium carbonate will no longer precipitate
magnesium carbonate. A mixture of solutions of magnesium and ammonium
chlorides, on evaporation or refrigeration, gives a double salt,
Mg(NH_{4})Cl_{3},6H_{2}O.[22] The salts of potassium, like those
of ammonium, are able to enter into combination with the magnesium
salts.[23] For instance, the double salt, MgKCl_{3},6H_{2}O, which is
known as _carnallite_,[24] and occurs in the salt mines of Stassfurt,
may be formed by freezing a saturated solution of potassium chloride
with an excess of magnesium chloride. A saturated solution of magnesium
sulphate dissolves potassium sulphate, and solid magnesium sulphate
is soluble in a saturated solution of potassium sulphate. A double
salt, K_{2}Mg(SO_{4})_{2},6H_{2}_O, which closely resembles the
above-mentioned ammonium salt, crystallises from these solutions.[25]
The nearest analogues of magnesium are able to give exactly similar
double salts, both in crystalline form (monoclinic system) and
composition; they, like this salt (_see_ Chapter XV.), are easily
able (at 140°) to part with all their water of crystallisation,
and correspond with the salts of sulphuric acid, whose type may
be taken as _magnesium sulphate_, MgSO_{4}.[26] It occurs at
Stassfurt as _kieserite_, MgSO_{4},H_{2}O, and generally separates
from solutions as a heptahydrated salt, MgSO_{4},7H_{2}O, and from
supersaturated solutions as a hexahydrated salt, MgSO_{4},6H_{2}O; at
temperatures below 0° it crystallises out as a dodecahydrated salt,
MgSO_{4},12H_{2}O, and a solution of the composition MgSO_{4},2H_{2}O
solidifies completely at -5°.[27] Thus between water and magnesium
sulphate there may exist several definite and more or less stable
degrees of equilibrium; the double salt MgSO_{4}K_{2}SO_{4},6H_{2}O
may be regarded as one of these equilibrated systems, the more so
since it contains 6H_{2}O, whilst MgSO_{4} forms its most stable
system with 7H_{2}O, and the double salt may be considered as this
crystallo-hydrate in which one molecule of water is replaced by the
molecule K_{2}SO_{4}.[28]
[20] Even a solution of ammonium chloride gives this salt with
magnesium sulphate. Its sp. gr. is 1·72; 100 parts of water at 0°
dissolve 9, at 20° 17·9 parts of the anhydrous salt. At about 130°
it loses all its water.
[21] This is an example of equilibrium and of the influence of mass;
the double salt is decomposed by water, but if instead of water we
take a solution of that soluble part which is formed in the
decomposition of the double salt, then the latter dissolves as a
whole.
[22] If an excess of ammonia be added to a solution of magnesium
chloride, only half the magnesium is thrown down in the
precipitate, 2MgCl_{2} + 2NH_{4}.OH = Mg(OH)_{2} + Mg.NH_{4}Cl_{3}
+ NH_{4}Cl. A solution of ammonium chloride reacts with magnesia,
evolving ammonia and forming a solution of the same salt, MgO +
3NH_{4}Cl = MgNH_{4}Cl_{3} + H_{2}O + 2NH_{3}.
Among the double salts of ammonium and magnesium, the phosphate,
MgNH_{4}PO_{4},6H_{2}O, is almost insoluble in water (0·07 gram is
soluble in a litre), even in the presence of ammonia. Magnesia is
very frequently precipitated as this salt from solutions in which
it is held by ammonium salts. As lime is not retained in solution
by the presence of ammonium salts, but is precipitated
nevertheless by sodium carbonate, &c., it is very easy to separate
calcium from magnesium by taking advantage of these properties.
[23] In order to see the nature and cause of formation of double salts,
it is sufficient (although this does not embrace the whole essence
of the matter) to consider that one of the metals of such salts
(for instance, potassium) easily gives acid salts, and the other
(in this instance, magnesium) basic salts; the properties of
distinctly basic elements predominate in the former, whilst in the
latter these properties are enfeebled, and the salts formed by
them bear the character of acids--for example, the salts of
aluminium or magnesium act in many cases like acids. By their
mutual combination these two opposite properties of the salts are
both satisfied.
[24] Carnallite has been mentioned in Chapter X. (Note 4) and in
Chapter XIII. These deposits also contain much _kainite_,
KMgCl(SO_{4}),3H_{2}O (sp. gr. 2·13; 100 parts of water dissolve
79·6 parts at 18°). This double salt contains two metals and two
haloids. Feit (1889) also obtained a bromide corresponding to
carnallite.
[25] The component parts of certain double salts diffuse at different
rates, and as the diffused solution contains a different
proportion of the component salts than the solution taken of the
double salt, it shows that such salts are decomposed by water.
According to Rüdorff, the double salts, like carnallite,
MgK_{2}(SO_{4})_{2},6H_{2}O, and the alums, all belong to this
order (1888). But such salts as tartar emetic, the double
oxalates, and double cyanides are not separated by diffusion,
which in all probability depends both on the relative rate of the
diffusion of the component salts and on the degree of affinity
acting between them. Those complex states of equilibrium which
exist between water, the individual salts MX and NY, and the
double salt MNXY, have been already partially analysed (as will be
shown hereafter) in that case when the system is heterogeneous
(that is, when something separates out in a solid state from the
liquid solution), but in the case of equilibria in a homogeneous
liquid medium (in a solution) the phenomenon is not so clear,
because it concerns that very theory of solution which cannot yet
be considered as established (Chapter I., Note 9, and others). As
regards the heterogeneous decomposition of double salts, it has
long been known that such salts as carnallite and
K_{2}Mg(SO_{4})_{2} give up the more soluble salt if an
insufficient quantity of water for their complete solution be
taken. The complete saturation of 100 parts of water requires at
0° 14·1, at 20° 25, and at 60° 50·2 parts of the latter double
salt (anhydrous), while 100 parts of water dissolve 27 parts of
magnesium sulphate at 0°, 36 parts at 20°, and 55 parts at 60°, of
the anhydrous salt taken. Of all the states of equilibrium
exhibited by double salts the most fully investigated as yet is
the system containing water, sodium sulphate, magnesium sulphate,
and their double salt, Na_{2}Mg(SO_{4})_{2}, which crystallises
with 4 and 6 mol. OH_{2}. The first crystallo-hydrate,
MgNa_{2}(SO_{4})_{2},4H_{2}O, occurs at Stassfurt, and as a
sedimentary deposit in many of the salt lakes near Astrakhan, and
is therefore called _astrakhanite_. The specific gravity of the
monoclinic prisms of this salt is 2·22. If this salt, in a finely
divided state, be mixed with the necessary quantity of water
(according to the equation MgNa_{2}(SO_{4})_{2},4H_{2}O + 13H_{2}O
= Na_{2}SO_{4},10H_{2}O + MgSO_{4},7H_{2}O), the mixture
solidifies like plaster of Paris into a homogeneous mass if the
temperature be _below_ 22° (Van't Hoff und Van Deventer, 1886;
Bakhuis Roozeboom, 1887); but if the temperature be above this
_transition-point_ the water and double salt do not react on each
other: that is, they do not solidify or give a mixture of sodium
and magnesium sulphates. If a mixture (in equivalent quantities)
of solutions of these salts be evaporated, and crystals of
astrakhanite and of the individual salts capable of proceeding
from it be added to the concentrated solution to avoid the
possibility of a supersaturated solution, then at temperatures
above 22° astrakhanite is exclusively formed (this is the method
of its production), but at lower temperatures the individual salts
are alone produced. If equivalent amounts of Glauber's salt and
magnesium sulphate be mixed together in a solid state, there is no
change at temperatures below 22°, but at higher temperatures
astrakhanite and water are formed. The volume occupied by
Na_{2}SO_{4},10H_{2}O in grams = 322/1·46 = 220·5 cubic
centimetres, and by MgSO_{4},7H_{2}O = 246/1·68 = 146·4 c.c.;
hence their mixture in equivalent quantities occupies a volume of
366·9 c.c. The volume of astrakhanite = 334/2·22 = 150·5 c.c., and
the volume of 13H_{2}O = 234 c.c., hence their sum = 380·5 c.c.,
and therefore it is easy to follow the formation of the
astrakhanite in a suitable apparatus (a kind of thermometer
containing oil and a powdered mixture of sodium and magnesium
sulphates), and to see by the variation in volume that below 22°
it remains unchanged, and at higher temperatures proceeds the more
quickly the higher the temperature. At the transition temperature
the solubility of astrakhanite and of the mixture of the component
salts is one and the same, whilst at higher temperatures a
solution which is saturated for a mixture of the individual salts
would be supersaturated for astrakhanite, and at lower
temperatures the solution of astrakhanite will be supersaturated
for the component salts, as has been shown with especial detail by
Karsten, Deacon, and others. Roozeboom showed that there are two
limits to the composition of the solutions which can exist for a
double salt; these limits are respectively obtained by dissolving
a mixture of the double salt with each of its component simple
salts. Van't Hoff demonstrated, besides this, that the tendency
towards the formation of double salts has a distinct influence on
the progress of double decomposition, for at temperatures above
31° the mixture 2MgSO_{4},7H_{2}O + 2NaCl passes into
MgNa_{2}(SO_{4})_{2},4H_{2}O + MgCl_{2},6H_{2}O + 4H_{2}O, whilst
below 31° there is not this double decomposition, but it proceeds
in the opposite direction, as may be demonstrated by the
above-described methods. Van der Heyd obtained a potassium
astrakhanite, K_{2}SO_{4}MgSO_{4},4H_{2}O, from solutions of the
component salts at 100°.
From these experiments on double salts we see that there is as
close a dependence between the temperature and the formation of
substances as there is between the temperature and a change of
state. It is a case of Deville's principles of dissociation,
extended in the direction of the passage of a solid into a liquid.
On the other hand, we see here how essential a _rôle_ water plays
in the formation of compounds, and how the affinity for water of
crystallisation is essentially analogous to the affinity between
salts, and hence also to the affinity of acids for bases, because
the formation of double salts does not differ in any essential
point (except the degree of affinity--that is, from a quantitative
aspect) from the formation of salts themselves. When sodium
hydroxide with nitric acid gives sodium nitrate and water the
phenomenon is essentially the same as in the formation of
astrakhanite from the salts Na_{2}SO_{4},10H_{2}O and
MgSO_{4},7H_{2}O. Water is disengaged in both cases, and hence the
volumes are altered.
[26] This salt, and especially its crystallo-hydrate with 7H_{2}O, is
generally known as Epsom salts. It has long been used as a
purgative. It is easily obtained from magnesia and sulphuric acid,
and it separates on the evaporation of sea water and of many
saline springs. When carbonic anhydride is obtained by the action
of sulphuric acid on magnesite, magnesium sulphate remains in
solution. When dolomite--that is, a mixture of magnesium and
calcium carbonates--is subjected to the action of a solution of
hydrochloric acid until about half of the salt remains, the
calcium carbonate is mostly dissolved and magnesium carbonate is
left, which by treatment with sulphuric acid gives a solution of
magnesium sulphate.
[27] The anhydrous salt, MgSO_{4} (sp. gr. 2·61), attracts moisture
(7 mol. H_{2}O) from moist air; when heated in steam or hydrogen
chloride it gives sulphuric acid, and when heated with carbon it
is decomposed according to the equation 2MgSO_{4} + C = 2SO_{2} +
CO_{2} + 2MgO. The monohydrated salt (kieserite), MgSO_{4},H_{2}O
(sp. gr. 2·56), dissolves in water with difficulty; it is formed
by heating the other crystallo-hydrates to 135°. The hexahydrated
salt is dimorphous. If a solution, saturated at the boiling-point,
be prepared, and cooled without access of crystals of the
heptahydrated salt, then MgSO_{4},6H_{2}O crystallises out in
_monoclinic_ prisms (Loewel, Marignac), which are quite as
unstable as the salt, Na_{2}SO_{4},7H_{2}O; but if prismatic
crystals of the cubic system of the copper-nickel salts of the
composition MSO_{4},6H_{2}O be added, then crystals of
MgSO_{4},6H_{2}O are deposited on them as prisms of the _cubic_
system (Lecoq de Boisbaudran). The common crystallo-hydrate,
MgSO_{4},7H_{2}O, Epsom salts, belongs to the _rhombic_ system,
and is obtained by crystallisation below 30°. Its specific gravity
is 1·69. In a vacuum, or at 100°, it loses 5H_{2}O, at 132°
6H_{2}O, and at 210° all the 7H_{2}O (Graham). If crystals of
ferrous or cobaltic sulphate be placed in a saturated solution,
_hexagonal_ crystals of the heptahydrated salt are formed (Lecoq
de Boisbaudran); they present an unstable state of equilibrium,
and soon become cloudy, probably owing to their transformation
into the more stable common form. Fritzsche, by cooling saturated
solutions below 0°, obtained a mixture of crystals of ice and of a
dodecahydrated salt, which easily split up at temperatures above
0°. Guthrie showed that dilute solutions of magnesium sulphate,
when refrigerated, separate ice until the solution attains a
composition MgSO_{4},24H_{2}O, which will completely freeze into a
crystallo-hydrate at -5·3°. According to Coppet and Rüdorff, the
temperature of the formation of ice falls by 0·073° for every part
by weight of the heptahydrated salt per 100 of water. This figure
gives (Chapter I., Note 49) _i_ = 1 for both the heptahydrated and
the anhydrous salt, from which it is evident that it is impossible
to judge the state of combination in which a dissolved substance
occurs by the temperature of the formation of ice.
The solubility of the different crystallo-hydrates of magnesium
sulphate, according to Loewel, also varies, like those of sodium
sulphate or carbonate (_see_ Chapter XII., Notes 7 and 18). At 0°
100 parts of water dissolves 40·75 MgSO_{4} in the presence of the
hexahydrated salt, 34·67 MgSO_{4} in the presence of the hexagonal
heptahydrated salt, and only 26 parts of MgSO_{4} in the presence
of the ordinary heptahydrated salt--that is, solutions giving the
remaining crystallo-hydrates will be supersaturated for the
ordinary heptahydrated salt.
All this shows how many diverse aspects of more or less stable
equilibria may exist between water and a substance dissolved in
it; this has already been enlarged on in Chapter I.
Carefully purified magnesium sulphate in its aqueous solution
gives, according to Stcherbakoff, an alkaline reaction with
litmus, and an acid reaction with phenolphthalein.
The specific gravity of solutions of certain salts of magnesium
and calcium reduced to 15°/4° (see my work cited, Chapter I., Note
19), are, if water at 4° = 10,000,
MgSO_{4}: _s_ = 9,992 + 99·89_p_ + 0·553_p_^2
MgCl_{2}: _s_ = 9,992 + 81·31_p_ + 0·372_p_^2
CaCl_{2}: _s_ = 9,992 + 80·24_p_ + 0·476_p_^2
[28] Graham even distinguished the last equivalent of the water of
crystallisation of the heptahydrated salt as that which is
replaced by other salts, pointing out that double salts like
MgK_{2}(SO_{4})_{2},6H_{2}O lose all their water at 135°, whilst
MgSO_{4},7H_{2}O only parts with 6H_{2}O.
_The power of forming basic salts_ is a very remarkable peculiarity of
magnesia and other feeble bases, and especially of those corresponding
with polyvalent metals. The very powerful bases corresponding with
univalent metals--like potassium and sodium--do not form basic salts,
and, indeed, are more prone to give acid salts, whilst magnesium easily
and frequently forms basic salts, especially with feeble acids, although
there are some oxides--as, for example, copper and lead oxides--which
still more frequently give basic salts. If a cold solution of magnesium
sulphate be mixed with a solution of sodium carbonate there is formed a
gelatinous precipitate of a basic salt, Mg(HO)_{2},4MgCO_{3},9H_{2}O;
but all the magnesia is not precipitated in this case, as a portion of it
remains in solution as an acid double salt. If sodium carbonate be added
to a boiling solution of magnesium sulphate a precipitate of a still more
basic salt is formed, 4MgSO_{4} + 4Na_{2}CO_{3} + 4H_{2}O = 4Na_{2}SO_{4}
+ CO_{2} + Mg(OH)_{2},3MgCO_{3},3H_{2}O. This basic salt forms the
ordinary drug _magnesia_ (_magnesia alba_), in the form of light porous
lumps. Other basic salts are formed under certain modifications of
temperature and conditions of decomposition. But _the normal salt_,
MgCO_{3}, which occurs in nature as magnesite in the form of rhombohedra
of specific gravity 3·056, cannot be obtained by such a method of
precipitation. In fact, the formation of the different basic salts shows
the power of water to decompose the normal salt. It is possible, however,
to obtain this salt both in an anhydrous and hydrated state. A solution
of magnesium carbonate in water containing carbonic acid is taken for
this purpose. The reason for this is easily understood--carbonic
anhydride is one of the products of the decomposition of magnesium
carbonate in the presence of water. If this solution be left to evaporate
spontaneously the normal salt separates in a hydrated form, but in the
evaporation of a heated solution, through which a stream of carbonic
anhydride is passed, the anhydrous salt is formed as a crystalline mass,
which remains unaltered in the air, like the natural mineral.[29] The
decomposing influence of water on the salts of magnesium, which is
directly dependent on the feeble basic properties of magnesia,[30] is
most clearly seen in _magnesium chloride_, MgCl_{2}. This salt is
contained[31] in the last mother-liquors of the evaporation of sea-water.
On cooling a sufficiently concentrated solution, the crystallo-hydrate,
MgCl_{2},6H_{2}O, separates;[32] but if it be further heated (above
106°) to remove the water, then hydrochloric acid passes off together
with the latter, so that there ultimately remains magnesia with a small
quantity of magnesium chloride.[33] From what has been said it is evident
that anhydrous magnesium chloride cannot be obtained by simple
evaporation. But if sal-ammoniac or sodium chloride be added to a
solution of magnesium chloride, then the evolution of hydrochloric acid
does not take place, and after complete evaporation the residue is
perfectly soluble in water. This renders it possible to obtain anhydrous
magnesium chloride from its aqueous solution. Indeed the mixture with
sal-ammoniac (in excess) may be dried (the residue consists of an
anhydrous double salt, MgCl_{2},2NH_{4}Cl) and then ignited (460°), when
the sal-ammoniac is converted into vapour and a fused mass of anhydrous
magnesium chloride remains behind. The anhydrous chloride evolves a very
considerable amount of heat on the addition of water, which shows the
great affinity the salt has for water.[34] Anhydrous magnesium chloride
is not only obtained by the above method, but is also formed by the
direct combination of chlorine and magnesium, and by the action of
chlorine on magnesium oxide, oxygen being evolved; this proceeds still
more easily _by heating magnesia with charcoal in a stream of chlorine_,
when the charcoal serves to take up the oxygen. This latter method is
also employed for the preparation of chlorides which are formed in an
anhydrous condition with still greater difficulty than magnesium
chloride. Anhydrous magnesium chloride forms a colourless, transparent
mass, composed of flexible crystalline plates of a pearly lustre. It
fuses at a low red heat (708°) into a colourless liquid, remains
unchanged in a dry state, but under the action of moisture is partially
decomposed even at the ordinary temperature, with formation of
hydrochloric acid. When heated in the presence of oxygen (air) it gives
chlorine and the basic salt, which is formed with even greater facility
under the action of heat in the presence of steam, when HCl is formed,
according to the equation 2MgCl_{2} + H_{2}O = MgOMgCl_{2} + 2HCl.[34 bis]
[29] The crystalline form of the anhydrous salt obtained in
this manner is not the same as that of the natural salt. The
former gives rhombohedra, like those in which calcium carbonate
appears as calc spar, whilst the natural salt appears as rhombic
prisms, like those sometimes presented by the same carbonate as
aragonite, which will shortly be described.
[30] Magnesium sulphate enters into certain reactions which are proper
to sulphuric acid itself. Thus, for instance, if a carefully
prepared mixture of equivalent quantities of hydrated magnesium
sulphate and sodium chloride be heated to redness, the evolution
of hydrochloric acid is observed just as in the action of
sulphuric acid on common salt, MgSO_{4} + 2NaCl + H_{2}O =
Na_{2}SO_{4} + MgO + 2HCl. Magnesium sulphate acts in a similar
manner on nitrates, with the evolution of nitric acid. A mixture
of it with common salt and manganese peroxide gives chlorine.
Sulphuric acid is sometimes replaced by magnesium sulphate in
galvanic batteries--for example, in the well-known Meidinger
battery. In the above-mentioned reactions we see a striking
example of the similarity of the reactions of acids and salts,
especially of salts which contain such feeble bases as magnesia.
[31] As sea-water contains many salts, MCl and MgX_{2}, it follows,
according to Berthollet's teaching, that MgCl_{2} is also present.
[32] As the crystallo-hydrates of the salts of sodium often contain
10H_{2}O, so many of the salts of magnesium contain 6H_{2}O.
[33] This decomposition is most simply defined as the result of the
two reverse reactions, MgCl_{2} ÷ H_{2}O = MgO + 2HCl and MgO +
2HCl = MgCl_{2} + H_{2}O, or as a distribution between O and
Cl_{2} on the one hand and H_{2} and Mg on the other. (With O,
MgCl_{2} gives chlorine, _see_ Chapter X., Note 33, and Chapter
II., Note 3 bis and others, where the reactions and
applications of MgCl_{2} are given.) It is then clear that,
according to Berthollet's doctrine, the mass of the hydrochloric
acid converts the magnesium oxide into chloride, and the mass of
the water converts the magnesium chloride into oxide. The
crystallo-hydrate, MgCl_{2},6H_{2}O, forms the limit of the
reversibility. But an intermediate state of equilibrium may exist
in the form of basic salts. On mixing ignited magnesia with a
solution of magnesium chloride of specific gravity about 1·2, a
solid mass is obtained which is scarcely decomposed by water at
the ordinary temperature (_see_ Chapter XVI., Note 4). A similar
means is employed for cementing sawdust into a solid mass, called
cylolite, used for flooring, &c.
We may remark that MgBr_{2} crystallises not only with 6H_{2}O
(temperature of fusion 152°), but also with 10H_{2}O (temperature
of fusion +12°, formed at -18°). (Panfiloff, 1894).
[34] According to Thomsen, the combination of MgCl_{2} with 6H_{2}O
evolves 33,000 calories, and its solution in an excess of water
36,000.
[34 bis] Hence MgCl_{2} may be employed for the preparation of chlorine
and hydrochloric acid (Chapters X. and XI.). In general magnesium
chloride, which is obtained in large quantities from sea water and
Stassfurt carnallite, may find numerous practical uses.
_Calcium_ (or the metal of lime) and its compounds in many respects
present a great resemblance to magnesium compounds, but are also clearly
distinguished from them by many properties.[35] In general, calcium
stands to magnesium in the same relation as potassium occupies in respect
to sodium. Davy obtained metallic calcium, like potassium, as an amalgam
by the action of a galvanic current; but neither charcoal nor iron
decomposes calcium oxide, and even sodium decomposes calcium chloride[36]
with difficulty. But a galvanic current easily decomposes calcium
chloride, and metallic sodium somewhat easily decomposes calcium iodide
when heated. As in the case of hydrogen, potassium, and magnesium, the
affinity of iodine for calcium is feebler than that of chlorine (and
oxygen), and therefore it is not surprising that calcium iodide may be
subjected to that decomposition, which the chloride and oxide undergo
with difficulty.[37] _Metallic calcium_ is of a yellow colour, and has a
considerable lustre, which it preserves in dry air. Its specific gravity
is 1·58. Calcium is distinguished by its great ductility; it melts at a
red heat and then burns in the air with a very brilliant flame; the
brilliancy is due to the formation of finely divided infusible calcium
oxide. Judging from the fact that calcium in burning gives a very large
flame, it is probable that this metal is volatile. Calcium decomposes
water at the ordinary temperature, and is oxidised in moist air, but not
so rapidly as sodium. In burning, it gives its oxide or _lime_, CaO, a
substance which is familiar to every one, and of which we have already
frequently had occasion to speak. This oxide is not met with in nature in
a free state, because it is an energetic base which everywhere encounters
acid substances forming salts with them. It is generally combined with
silica, or occurs as calcium carbonate or sulphate. The carbonate and
nitrate are decomposed, at a red heat, with the formation of lime. As a
rule, the carbonate, which is so frequently met with in nature, serves as
the source of the calcium oxide, both commercial and pure. When heated,
calcium carbonate dissociates: CaCO_{3} = CaO + CO_{2}. In practice the
decomposition is conducted at a bright red heat, in the presence of
steam, or a current of a foreign gas, in heaps or in special kilns.[38]
[35] There are many other methods of separating calcium from magnesium
besides that mentioned above (Note 22). Among them it will be
sufficient to mention the behaviour of these bases towards a
solution of sugar; hydrated _lime_ is exceedingly _soluble in an
aqueous solution of sugar_, whilst magnesia is but little soluble.
All the lime may be extracted from dolomite by burning it, slaking
the mixture of oxides thus obtained, and adding a 10 p.c. solution
of sugar. Carbonic anhydride precipitates calcium carbonate from
this solution. The addition of sugar (molasses) to the lime used
for building purposes powerfully increases the binding power of
the mortar, as I have myself found. I have been told that in the
East (India, Japan) the addition of sugar to cement has long been
practised.
[36] Moreover Caron obtained an alloy of calcium and zinc by fusing
calcium chloride with zinc and sodium. The zinc distilled from
this alloy at a white heat, leaving calcium behind (Note 50).
[37] Calcium iodide may be prepared by saturating lime with hydriodic
acid. It is a very soluble salt (at 20° one part of the salt
requires 0·49 part and at 43° 0·35 part of water for solution), is
deliquescent in the air, and resembles calcium chloride in many
respects. It changes but little when evaporated, and like calcium
chloride fuses when heated, and therefore all the water may be
driven off by heat. If anhydrous calcium iodide be heated with an
equivalent quantity of sodium in a closely covered iron crucible,
sodium iodide and metallic calcium are formed (Liés-Bodart). Dumas
advises carrying on this reaction in a closed space under
pressure.
[38] Kilns which act either intermittently or continuously are built
for this purpose. Those of the first kind are filled with
alternate layers of fuel and limestone; the fuel is lighted, and
the heat developed by its combustion serves for decomposing the
limestone. When the process is completed the kiln is allowed to
cool somewhat, the lime raked out, and the same process repeated.
In the continuously acting furnaces, constructed like that shown
in fig. 78, the kiln itself only contains limestone, and there are
lateral hearths for burning the fuel, whose flame passes through
the limestone and serves for its decomposition. Such furnaces are
able to work continuously, because the unburnt limestone may be
charged from above and the burnt lime raked out from below. It is
not every limestone that is suitable for the preparation of lime,
because many contain impurities, principally clay, dolomite, and
sand. Such limestones when burnt either fuse partially or give an
impure lime, called _poor_ lime in distinction from that obtained
from purer limestone, which is called _rich_ lime. The latter kind
is characterised by its disintegrating into a fine powder when
treated with water, and is suitable for the majority of uses to
which lime is applied, and for which the poor lime is sometimes
quite unfit. However, certain kinds of poor lime (as we shall see
in Chapter XVIII., Note 25) are used in the preparation of
hydraulic cements, which solidify into a hard mass under water.
In order to obtain perfectly pure lime it is necessary to take the
purest possible materials. In the laboratory, marble or shells are
used for this purpose as a pure form of calcium carbonate. They
are first burnt in a furnace, then put in a crucible and moistened
with a small quantity of water, and finally strongly ignited, by
which means a pure lime is obtained. Pure lime may be more rapidly
prepared by taking calcium nitrate, CaN_{2}O_{6}, which is easily
obtained by dissolving limestone in nitric acid. The solution
obtained is boiled with a small quantity of lime in order to
precipitate the foreign oxides which are insoluble in water. The
oxides of iron, aluminium, &c., are precipitated by this means.
The salt is then crystallised and ignited: CaN_{2}O_{6} = CaO +
2NO_{2} + O.
In the decomposition of calcium carbonate the lime preserves the
form of the lumps subjected to ignition; this is one of the signs
distinguishing quicklime when it is freshly burnt and unaltered by
air. It attracts moisture from the air and then disintegrates to a
powder; if left long exposed in the air, it also attracts carbonic
anhydride and increases in volume; it does not entirely pass into
carbonate, but forms a compound of the latter with caustic lime.
[Illustration: FIG. 78.--Continually-acting kiln for burning lime. The
lime is charged from above and calcined by four lateral grates, R, M. D,
fire-bars. B, space for withdrawing the burnt lime. K, stoke-house. M.
fire grate. Q, R, under-grate.]
Calcium oxide--that is, quicklime--is a substance (sp. gr. 3·15) which
is unaffected by heat,[39] and may therefore serve as a fire-resisting
material, and was employed by Deville for the construction of furnaces in
which platinum was melted, and silver volatilised by the action of the
heat evolved by the combustion of detonating gas. The hydrated lime,
slaked lime, or calcium hydroxide, CaH_{2}O_{2} (specific gravity 2·07)
is a most common alkaline substance, employed largely in building for
making mortars or cements, in which case its binding property is mainly
due to the absorption of carbonic anhydride.[40] Lime, like other
alkalis, acts on many animal and vegetable substances, and for this
reason has many practical uses--for example, for removing fats, and in
agriculture for accelerating the decomposition of organic substances in
the so-called _composts_ or accumulations of vegetable and animal remains
used for fertilising land. Calcium hydroxide easily loses its water at a
moderate heat (530°), but it does not part with water at 100°. When mixed
with water, lime forms a pasty mass known as _slaked lime_ and in a more
dilute form as _milk of lime_, because when shaken up in water it remains
suspended in it for a long time and presents the appearance of a milky
liquid. But, besides this, lime is directly soluble in water, not to any
considerable extent, but still in such a quantity that _lime water_ is
precipitated by carbonic anhydride, and has clearly distinguishable
alkaline properties. One part of lime requires at the ordinary
temperature about 800 parts of water for solution. At 100° it requires
about 1500 parts of water, and therefore lime-water becomes cloudy when
boiled. If lime-water be evaporated in a vacuum, calcium hydroxide
separates in six-sided crystals.[41] If lime-water be mixed with hydrogen
peroxide minute crystals of _calcium peroxide_, CaO_{2},8H_{2}O,
separate; this compound is very unstable and, like barium peroxide, is
decomposed by heat. Lime, as a powerful base, combines with all acids,
and in this respect presents a transition from the true alkalis to
magnesia. Many of the salts of calcium (the carbonate, phosphate,
borate, and oxalate) are insoluble in water; besides which the sulphate
is only sparingly soluble. As a more energetic base than magnesia, lime
forms salts, CaX_{2}, which are distinguished by their stability in
comparison with the salts MgX_{2}; neither does lime so easily form basic
and double salts as magnesia.
[39] Lime, when raised to a white heat in the vapour of potassium,
gives calcium, and in chlorine it gives off oxygen. Sulphur,
phosphorus, &c., when heated with lime, are absorbed by it.
[40] The greater quantity of lime is used in making mortar for binding
bricks or stones together, in the form of _lime_ or _cement_, or
the so-called _slaked lime_. For this purpose the lime is mixed
with water and sand, which serves to separate the particles of
lime from each other. If only lime paste were put between two
bricks they would not hold firmly together, because after the
water had evaporated the lime would occupy a smaller space than
before, and therefore cracks and powder would form in its mass, so
that it would not at all produce that complete cementation of the
bricks which it is desired to attain. Pieces of stone--that is,
sand--mixed with the lime hinder this process of disintegration,
because the lime binds together the individual grains of sand
mixed with it, and forms one concrete mass, in consequence of a
process which proceeds after the desiccation or removal of the
water. The process of the solidification of lime, taken as slaked
lime, consists first in the direct evaporation of the water and
crystallisation of the hydrate, so that the lime binds the stones
and sand mixed with it, just as glue binds two pieces of wood. But
this preliminary binding action of lime is feeble (as is seen by
direct experiment) unless there be further alteration of the lime
leading to the formation of carbonates, silicates, and other salts
of calcium which are distinguished by their great cohesiveness.
With the progress of time the cement is partially subjected to the
action of the carbonic anhydride in the air, owing to which
calcium carbonate is formed, but not more than half the lime is
thus converted into carbonate. Besides which, the lime partially
acts on the silica of the bricks, and it is owing to these new
combinations simultaneously forming in the cement that it
gradually becomes stronger and stronger. Hence the binding action
of the lime becomes stronger with the lapse of time. This is the
reason (and not, as is sometimes said, because the ancients knew
how to build stronger than we do) why buildings which have stood
for centuries possess a very strongly binding cement. Hydraulic
cements will be described later (Chapter XVIII., Note 25).
[41] Professor Glinka measured the transparent bright crystals of
calcium hydroxide which are formed in common hydraulic (Portland)
cement.
Anhydrous lime does not absorb dry carbonic anhydride at the ordinary
temperature. This was already known by Scheele, and Prof. Schuliachenko
showed that there is no absorption even at 360°. It only proceeds at a
red heat,[42] and then only leads to the formation of a mixture of
calcium oxide and carbonate (Rose). But if the lime be slaked or
dissolved, the absorption of carbonic anhydride proceeds rapidly and
completely. These phenomena are connected with the _dissociation of
calcium carbonate_, studied by Debray (1867) under the influence of the
conceptions of dissociation introduced into science by Henri Saint-Claire
Deville. Just as there is no vapour tension for non-volatile substances,
so there is no dissociation tension of carbonic anhydride for calcium
carbonate at the ordinary temperature. Just as every volatile substance
has a maximum possible vapour tension for every temperature, so also
calcium carbonate has its corresponding _dissociation tension_; this at
770° (the boiling point of cadmium) is about 85 mm. (of the mercury
column), and at 930° (the boiling point of Zn) it is about 520 mm. As, if
the tension be greater, there will be no evaporation, so also there will
he no decomposition. Debray took crystals of calc spar, and could not
observe the least change in them at the boiling point of zinc (930°) in
an atmosphere of carbonic anhydride taken at the atmospheric pressure
(760 mm.), whilst on the other hand calcium carbonate may be completely
decomposed at a much lower temperature if the tension of the carbonic
anhydride be kept below the dissociation tension, which may be done
either by directly pumping away the gas with an air-pump, or by mixing it
with some other gas--that is, by diminishing the partial pressure of the
carbonic anhydride,[43] just as an object may be dried at the ordinary
temperature by removing the aqueous vapour or by carrying it off in a
stream of another gas. Thus it is possible to obtain calcium carbonate
from lime and carbonic anhydride at a certain temperature above that at
which dissociation begins, and conversely to decompose calcium carbonate
at the same temperature into lime and carbonic anhydride.[44] At the
ordinary temperature the reaction of the first order (combination) cannot
proceed because the second (decomposition, dissociation) cannot take
place, and thus all the most important phenomena with respect to the
behaviour of lime towards carbonic anhydride are explained by starting
from one common basis.[45]
[42] The act of heating brings the substance into that state of
internal motion which is required for reaction. It should be
considered that by the act of heating not only is the bond between
the parts, or cohesion of the molecules, altered (generally
diminished), not only is the motion or store of energy of the
whole molecule increased, but also that in all probability the
motion of the atoms themselves in molecules undergoes a change.
The same kind of change is accomplished by the act of solution, or
of combination in general, judging from the fact that a dissolved
or combined substance--for instance, lime with water--reacts on
carbonic anhydride as it does under the action of heat. For the
comprehension of chemical phenomena it is exceedingly useful to
recognise clearly this parallelism. Rose's observation on the
formation (by the slow diffusion of solutions of calcium chloride
and sodium carbonate) of aragonite from dilute, and of calc spar
from strong, solutions is easily understood from this point of
view. As aragonite is always formed from hot solutions, it appears
that dilution with water acts like heat. The following experiment
of Kühlmann is particularly instructive in this sense. Anhydrous
(perfectly dry) barium oxide does not react with monohydrated
sulphuric acid, H_{2}SO_{4} (containing neither free water nor
anhydride, SO_{3}). But if either an incandescent object or a
moist substance is brought into contact with the mixture a violent
reaction immediately begins (it is essentially the same as
combustion), and the whole mass reacts.
The influence of solution on the process of reaction is
instructively illustrated by the following experiment. Lime, or
barium oxide, is placed in a flask or retort having an upper
orifice and connected with a tube immersed in mercury. A funnel
furnished with a stopcock and filled with water is fixed into the
upper orifice of the retort, which is then filled with dry
carbonic anhydride. There is no absorption. When a constant
temperature is arrived at, the unslaked oxide is made to absorb
all the carbonic anhydride by carefully admitting water. A vacuum
is formed, as is seen by the mercury rising in the neck of the
retort. With water the absorption goes on to the end, whilst under
the action of heat there remains the dissociating tension of the
carbonic anhydride. Furthermore, we here see that, with a certain
resemblance, there is also a distinction, depending on the fact
that at low temperatures calcium carbonate does not dissociate;
this determines the complete absorption of the carbonic anhydride
in the aqueous solution.
[43] Experience has shown that by moistening partially-burnt lime with
water and reheating it, it is easy to drive off the last traces of
carbonic anhydride from it, and that, in general, by blowing air
or steam through the lime, and even by using moist fuel, it is
possible to accelerate the decomposition of the calcium carbonate.
The partial pressure is decreased by these means.
[44] Before the introduction of Deville's theory of dissociation, the
_modus operandi_ of decompositions like that under consideration
was understood in the sense that decomposition starts at a certain
temperature, and that it is accelerated by a rise of temperature,
but it was not considered possible that combination could proceed
at the same temperature as that at which decomposition goes on.
Berthollet and Deville introduced the conception of equilibrium
into chemical science, and elucidated the question of reversible
reactions. Naturally the subject is still far from being
clear--the questions of the rate and completeness of reaction, of
contact, &c., still intrude themselves--but an important step has
been made in chemical mechanics, and we have started on a new path
which promises further progress, towards which much has been done
not only by Deville himself, but more especially by the French
chemists Debray, Troost, Lemoine, Hautefeuille, Le Chatelier, and
others. Among other things those investigators have shown the
close resemblance between the phenomena of evaporation and
dissociation, and pointed out that the amount of heat absorbed by
a dissociating substance may be calculated according to the law of
the variation of dissociation-pressure, in exactly the same manner
as it is possible to calculate the latent heat of the evaporation
of water, knowing the variation of the tension with the
temperature, on the basis of the second law of the mechanical
theory of heat. Details of this subject must be looked for in
special works on physical chemistry. _One and the same conception_
of the mechanical theory of heat _is applicable to dissociation_
and _evaporation_.
[45] But the question as to the formation of a basic calcium carbonate
with a rise of temperature still remains undecided. The presence
of water complicates all the relations between lime and carbonic
anhydride, all the more as the existence of an attraction between
calcium carbonate and water is seen from its being able to give a
_crystallo-hydrate_, CaCO_{3},5H_{2}O (Pelouze), which
crystallises in rhombic prisms of sp. gr. about 1·77 and loses its
water at 20°. These crystals are obtained when a solution of lime
in sugar and water is left long exposed to the air and slowly
attracts carbonic anhydride from it, and also by the evaporation
of such a solution at a temperature of about 3°. On the other
band, it is probable that an _acid salt_, CaH_{2}(CO_{3})_{2}, is
formed in an aqueous solution, not only because water containing
carbonic acid dissolves calcium carbonate, but more especially in
view of the researches of Schloesing (1872), which showed that at
16° a litre of water in an atmosphere of carbonic anhydride
(pressure 0·984 atmosphere) dissolves 1·086 gram of calcium
carbonate and 1·778 gram of carbonic anhydride, which corresponds
with the formation of calcium hydrogen carbonate, and the solution
of carbonic anhydride in the remaining water. Caro showed that a
litre of water is able to dissolve as much as 3 grams of calcium
carbonate if the pressure be increased to 4 and more atmospheres.
The calcium carbonate is precipitated when the carbonic anhydride
passes off in the air or in a current of another gas; this also
takes place in many natural springs. Tufa, stalactites, and other
like formations from waters containing calcium carbonate and
carbonic acid in solution are formed in this manner. The
solubility of calcium carbonate itself at the ordinary temperature
does not exceed 13 milligrams per litre of water.
_Calcium carbonate_, CaCO_{3}, is sometimes met with in nature in a
crystalline form, and it forms an example of the phenomenon termed
_dimorphism_--that is, it appears in two crystalline forms. When it
exhibits combinations of forms belonging to the hexagonal system
(six-sided prisms, rhombohedra, &c.) it is called _calc spar_. Calc spar
has a specific gravity of 2·7, and is further characterised by a distinct
cleavage along the planes of the fundamental rhombohedron having an angle
of 105°. Perfectly transparent Iceland spar presents a clear example of
double refraction (for which reason it is frequently employed in physical
apparatus). The other form of calcium carbonate occurs in crystals
belonging to the rhombic system, and it is then called _aragonite_; its
specific gravity is 3·0. If calcium carbonate be artificially produced by
slow crystallisation at the ordinary temperature, it appears in the
rhombohedral form, but if the crystallisation be aided by heat it then
appears as aragonite. It may therefore be supposed that calc spar
presents the form corresponding with a low temperature, and aragonite
with a higher temperature during crystallisation.[46]
[46] Dimorphous bodies differ from true isomers and polymers in that
they do not differ in their chemical reactions, which are
determined by a difference in the distribution (motion) of the
atoms in the molecules, and therefore dimorphism is usually
ascribed to a difference in the distribution of similar molecules,
building up a crystal. Although such a hypothesis is quite
admissible in the spirit of the atomic and molecular theory, yet,
as in such a redistribution of the molecules a perfect
conservation of the distribution of the atoms in them cannot be
imagined, and in every effort of chemical reaction there must take
place a certain motion among the atoms; so in my opinion there is
no firm basis for distinguishing dimorphism from the general
conception of isomerism, under which the cases of those organic
bodies which are dextro and lævo rotatory (with respect to
polarised light) have recently been brought with such brilliant
success. When calcium carbonate separates out from solutions, it
has at first a gelatinous appearance, which leads to the
supposition that this salt appears in a colloidal state. It only
crystallises with the progress of time. The colloidal state of
calcium carbonate is particularly clear from the following
observations made by Prof. Famintzin, who showed that when it
separates from solutions it is obtained under certain conditions
in the form of grains having the peculiar paste-like structure
proper to starch, which fact has not only an independent interest,
but presents an example of a mineral substance being obtained in a
form until then only known in the organic substances elaborated in
plants. This shows that the forms (cells, vessels, &c.) in which
vegetable and animal substances occur in organisms do not present
in themselves anything peculiar to organisms, but are only the
result of those particular conditions in which these substances
are formed. Traube and afterwards Monnier and Vogt (1882) obtained
formations which, under the microscope, were in every respect
identical in appearance with vegetable cells, by means of a
similar slow formation of precipitates (by reacting on sulphates
of different metals with sodium silicate or carbonate).
_Calcium sulphate_ in combination with two equivalents of water,
CaSO_{4},2H_{2}O, is very widely distributed in nature, and is known as
_gypsum_. Gypsum loses one and a half and two equivalents of water at a
moderate temperature,[47] and anhydrous or burnt gypsum is then obtained,
which is also known as plaster of Paris, and is employed in large
quantities for modelling.[48] This use depends on the fact that burnt and
finely-divided and sifted gypsum forms a paste when mixed with water;
after a certain time this paste becomes slightly heated and solidifies,
owing to the fact that the anhydrous calcium sulphate, CaSO_{4}, again
combines with water. When the plaster of Paris and water are first made
into a paste they form a mechanical mixture, but when the mass
solidifies, then a compound of the calcium sulphate with two molecules of
water is produced; and this may be regarded as derived from S(OH)_{6} by
the substitution of two atoms of hydrogen by one atom of bivalent
calcium. Natural gypsum sometimes appears as perfectly colourless, or
variegated, marble-like, masses, and sometimes in perfectly colourless
crystals, _selenite_, of specific gravity 2·33. The semi-transparent
gypsum, or _alabaster_, is often carved into small statues. Besides
which an anhydrous calcium sulphate, CaSO_{4}, called _anhydrite_
(specific gravity 2·97), occurs in nature. It sometimes occurs along with
gypsum. It is no longer capable of combining directly with water, and
differs in this respect from the anhydrous salt obtained by gently
igniting gypsum. If gypsum be very strongly heated it shrinks and loses
its power of combining with water.[48 bis] One part of calcium sulphate
requires at 0° 525 parts of water for solution, at 38° 466 parts, and at
100° 571 parts of water. The maximum solubility of gypsum is at about
36°, which is nearly the same temperature as that at which sodium
sulphate is most soluble.[49]
[47] According to Le Chatelier (1888), 1-1/2H_{2}O is lost at
120°--that is, H_{2}O,2CaSO_{4} is formed, but at 194° all the
water is expelled. According to Shenstone and Cundall (1888)
gypsum begins to lose water at 70° in dry air. The semi-hydrated
compound H_{2}O,2CaSO_{4} is also formed when gypsum is heated
with water in a closed vessel at 150° (Hoppe-Seyler).
[48] For stucco-work it is usual to add lime and sand, as the mass is
then harder and does not solidify so quickly. For imitating
marble, glue is added to the plaster, and the mass is polished
when thoroughly dry. Re-burnt gypsum cannot be used over again, as
that which has once solidified is, like the natural anhydride, not
able to recombine with water. It is evident that the structure of
the molecules in the crystallised mass, or in general in any dense
mass, exerts an influence on the chemical action, which is more
particularly evident in metals in their different forms (powder,
crystalline, rolled, &c.)
[48 bis] According to MacColeb, gypsum dehydrated at 200° has a
specific gravity 2·577, and heated to its point of fusion, 2·654.
Potilitzin (1894) also admits the two above-named modifications of
anhydrous gypsum, which, moreover, always contain the
semi-hydrated hydrate (Note 47), and he explains by their relation
to water the phenomena observed in the solidification of a mixture
of burnt gypsum and water.
[49] As Marignac showed, gypsum, especially when desiccated at 120°,
easily gives supersaturated solutions with respect to
CaSO_{4},2H_{2}O, which contain as much as 1 part of CaSO_{4} to
110 parts of water. Boiling dilute hydrochloric acid dissolves
gypsum, forming calcium chloride. The behaviour of gypsum towards
the alkaline carbonates has been described in Chapter X. Alcohol
precipitates gypsum from its aqueous solutions, because, like the
sulphates in general, it is sparingly soluble in alcohol. Gypsum,
like all the sulphates, when heated with charcoal, gives up its
oxygen, forming the sulphide, CaS.
Calcium ulphate, like magnesium sulphate, is capable of forming
double salts, but with difficulty, and they are chemically less
stable. They contain, as is always the case with double salts,
less water of crystallisation than the component salts. Rose,
Struvé, and others obtained the salt CaK_{2}(SO_{4})_{2},H_{2}O; a
mixture of gypsum with an equivalent amount of potassium sulphate
and water solidifies into a homogeneous mass. Fritzsche obtained
the corresponding sodium salt in a hydrated and anhydrous state,
by heating a mixture of gypsum with a saturated solution of sodium
sulphate. The anhydrous salt occurs in nature as _glauberite_.
Fritzsche also obtained _gaylussite_,
Na_{2}Ca(CO_{3})_{2},5H_{2}O, by pouring a saturated solution of
sodium carbonate on to freshly-precipitated calcium carbonate.
Calcium also forms basic salts, but only a few. Veeren (1892)
obtained Ca(NO_{3})_{2}Ca(OH)_{2},2-1/2H_{2}O by leaving powdered
caustic lime in a saturated solution of Ca(NO_{3})_{2} until it
solidified. This salt is decomposed by water.
As lime is a more energetic base than magnesia, so _calcium chloride_,
CaCl_{2}, is not so easily decomposed by water, and its solutions only
disengage a small quantity of hydrochloric acid when evaporated, and when
the evaporation is conducted in a stream of hydrochloric acid it easily
gives an anhydrous salt which fuses at 719°; otherwise an aqueous
solution yields a crystallo-hydrate, CaCl_{2},6H_{2}O, which melts at
30°.[50]
[50] Calcium chloride has a specific gravity 2·20, or, when fused,
2·12, and the sp. gr. of the crystallised salt CaCl_{2},6H_{2}O is
1·69. If the volume of the crystals at 0° = 1, then at 29° it is
1·020, and the volume of the fused mass at the same temperature is
1·118 (Kopp) (specific gravity of solutions, _see_ Note 27). The
solution containing 50 p.c. CaCl_{2} boils at 130°, 70 p.c. at
158°. Superheated steam decomposes calcium chloride with more
difficulty than magnesium chloride and with greater ease than
barium chloride (Kuhnheim). Sodium does not decompose fused
calcium chloride even on prolonged heating (Liés-Bodart), but an
alloy of sodium with zinc, lead, and bismuth decomposes it,
forming an alloy of calcium with one of the above-named metals
(Caron). The zinc alloy may be obtained with as much as 15 p.c. of
calcium. Calcium chloride is soluble in alcohol and absorbs
ammonia.
A gram molecular weight of calcium chloride in dissolving in an
excess of water evolves 18,723 calories, and in dissolving in
alcohol 17,555 units of heat, according to Pickering.
Roozeboom made detailed researches on the crystallo-hydrates of
calcium chloride (1889), and found that CaCl_{2},6H_{2}O melts at
30·2°, and is formed at low temperatures from solutions containing
not more than 103 parts of calcium chloride per 100 parts of
water; if the amount of salt (always to 100 parts of water)
reaches 120 parts, then tabular crystals of
CaCl_{2},4H_{2}O[Greek: b] are formed, which at temperatures above
38·4° are converted into the crystallo-hydrates CaCl_{2},2H_{2}O,
whilst at temperatures below 18° the [Greek: b] variety passes
into the more stable CaCl_{2},4H_{2}O[Greek: a], which process is
aided by mechanical friction. Hence, as is the case with magnesium
sulphate (Note 27), one and the same crystallo-hydrate appears in
two forms--the [Greek: b], which is easily produced but is
unstable, and the [Greek: a], which is stable. The solubility of
the above-mentioned hydrates of chloride of calcium, or amount of
calcium chloride per 100 parts of water, is as follows:--
0° 20° 30° 40° 60°
CaCl_{2},6H_{2}O 60 75 100 (102·8)
CaCl_{2},4H_{2}O[Greek: a] -- 90 101 117} (154·2)
CaCl_{2},4H_{2}O[Greek: b] -- 104 114 --}
CaCl_{2},2H_{2}O -- -- (308·3) 128 137
The amount of calcium chloride to 100 parts of water in the
crystallo-hydrate is given in brackets. The point of intersection
of the curves of solubility lies at about 30° for the first two
salts and about 45° for the salts with 4H_{2}O and 2H_{2}O. The
crystals CaCl_{2},2H_{2}O may, however, be obtained (Ditte) at the
ordinary temperature from solutions containing hydrochloric acid.
The vapour tension of this crystallo-hydrate equals the
atmospheric at 165°, and therefore the crystals may be dried in an
atmosphere of steam and obtained without a mother liquor, whose
vapour tension is greater. This crystallo-hydrate decomposes at
about 175° into CaCl_{2},H_{2}O and a solution; this is easily
brought about in a closed vessel when the pressure is greater than
the atmosphere. This crystallo-hydrate is destroyed at
temperatures above 260°, anhydrous calcium chloride being formed.
Neglecting the unstable modification CaCl_{2},4H_{2}O[Greek: b],
we will give the temperatures _t_ at which the passage of one
hydrate into another takes place and at which the solution
CaCl_{2} + _n_H_{2}O, the two solids A and B and aqueous vapour,
whose tension is given as _p_ in millimetres, are able to exist
together in stable equilibrium, according to Roozeboom's
determinations:
_t_ _n_ A B _p_
-55° 14·5 ice CaCl_{2},6H_{2}O 0
+29·8° 6·1 CaCl_{2},6H_{2}O CaCl_{2},4H_{2}O 6·8
45·3° 4·7 CaCl_{2},4H_{2}O CaCl_{2},2H_{2}O 11·8
175·5° 2·1 CaCl_{2},2H_{2}O CaCl_{2},H_{2}O 842
260° 1·8 CaCl_{2},H_{2}O CaCl_{2} Several
atmospheres
Solutions of calcium chloride may serve as a convenient example
for the study of the supersaturated state, which in this case
easily occurs, because different hydrates are formed. Thus at 25°
solutions containing more than 83 parts of anhydrous calcium
chloride per 100 of water will be supersaturated for the hydrate
CaCl_{2},6H_{2}O.
On the other hand, Hammerl showed that solutions of calcium
chloride, when frozen, deposit ice if they contain less than 43
parts of salt per 100 of water, and if more the crystallo-hydrate
CaCl_{2},6H_{2}O separates, and that a solution of the above
composition (CaCl_{2},14H_{2}O requires 44·0 parts calcium
chloride per 100 of water) solidifies as a cryohydrate at
about -55°.
Just as for potassium, K = 39 (and sodium, Na = 23), there are the near
analogues, Rb = 85 and Cs = 133, and also another, Li = 7, so in exactly
the same manner for calcium, Ca = 40 (and magnesium, Mg = 24), there is
another analogue of lighter atomic weight, beryllium, Be = 9, besides the
near analogues strontium, Sr = 87, and barium, Ba = 137. As rubidium and
cæsium are more rarely met with in nature than potassium, so also
strontium and barium are rarer than calcium (in the same way that bromine
and iodine are rarer than chlorine). Since they exhibit many points of
resemblance with calcium, strontium and barium may be characterised after
a very short acquaintance with their chief compounds; this shows the
important advantages gained by distributing the elements according to
their natural groups, to which matter we shall turn our attention in the
next chapter.
Among the compounds of barium met with in nature the commonest is the
_sulphate_, BaSO_{4}, which forms anhydrous crystals of the rhombic
system, which are identical in their crystalline form with anhydrite,
and generally occur as transparent and semi-transparent masses of
tabular crystals having a high specific gravity, namely 4·45, for which
reason this salt bears the name of _heavy spar_ or _barytes_. Analogous
to it is _celestine_, SrSO_{4}, which is, however, more rarely met with.
Heavy spar frequently forms the gangue separated on dressing metallic
ores from the vein stuff; this mineral is the source of all other barium
compounds; for the carbonate, although more easily transformed into the
other compounds (because acids act directly on it, evolving carbonic
anhydride), is a comparatively rare mineral (BaCO_{3} forms the mineral
_witherite_; SrCO_{3}, _strontianite_; both are rare, the latter is
found at Etna). The treatment of barium sulphate is rendered difficult
from the fact that it is insoluble both in water and acids, and has
therefore to be treated by a method of reduction.[51] Like sodium
sulphate and calcium sulphate, heavy spar when heated with charcoal
parts with its oxygen and forms barium sulphide, BaS. For this purpose a
pasty mixture of powdered heavy spar, charcoal, and tar is subjected to
the action of a strong heat, when BaSO_{4} + 4C = BaS + 4CO. The residue
is then treated with water, in which the barium sulphide is soluble.[52]
When boiled with hydrochloric acid, barium chloride, BaCl_{2},
is obtained in solution, and the sulphur is disengaged as gaseous
sulphuretted hydrogen, BaS + 2HCl = BaCl_{2} + H_{2}S. In this manner
barium sulphate is converted into barium chloride,[53] and the latter
by double decomposition with strong nitric acid or nitre gives the less
soluble barium nitrate, Ba(NO_{3})_{2},[54] or with sodium carbonate a
precipitate of barium carbonate, BaCO_{3}. Both these salts are able to
give _barium oxide_, or _baryta_, BaO, and the hydroxide, Ba(HO)_{2},
which differs from lime by its great solubility in water,[55] and by the
ease with which it forms a crystallo-hydrate, BaH_{2}O_{2},8H_{2}O, from
its solutions. Owing to its solubility, baryta is frequently employed
in manufactures and in practical chemistry as an alkali which has the
very important property that it may be always entirely removed from
solution by the addition of sulphuric acid, which entirely separates
it as the insoluble barium sulphate, BaSO_{4}. It may also be removed
whilst it remains in an alkaline state (for example, the excess which
may remain when it is used for saturating acids) by means of carbonic
anhydride, which also completely precipitates baryta as a sparingly
soluble, colourless, and powdery carbonate. Both these reactions show
that baryta has such properties as would very greatly extend its use
were its compounds as widely distributed as those of sodium and calcium,
and were its soluble compounds not poisonous. Barium nitrate is directly
decomposed by the action of heat, barium oxide being left behind. The
same takes place with barium carbonate, especially that form of it
precipitated from solutions, and when mixed with charcoal or ignited
in an atmosphere of steam. Barium oxide combines with water with the
development of a large amount of heat, and the resultant hydroxide is
very stable in its retention of the water, although it parts with it
when strongly ignited.[55 bis] With oxygen the anhydrous oxide gives, as
already mentioned in Chapters III. and IV., a _peroxide_, BaO_{2}.[56]
Neither calcium nor strontium oxides are able to give such a peroxide
directly, but they form peroxides under the action of hydrogen peroxide.
[51] The action of barium sulphate on sodium and potassium carbonates
is given on p. 437.
[52] Barium sulphide is decomposed by water, BaS + 2H_{2}O = H_{2}S
+ Ba(OH)_{2} (the reaction is reversible), but both substances are
soluble in water, and their separation is complicated by the fact
that barium sulphide absorbs oxygen and gives insoluble barium
sulphate. The hydrogen sulphide is sometimes removed from the
solution by boiling with the oxides of copper or zinc. If sugar be
added to a solution of barium sulphide, barium saccharate is
precipitated on heating; it is decomposed by carbonic anhydride,
so that barium carbonate is formed. An equivalent mixture of
sodium sulphate with barium or strontium sulphates when ignited
with charcoal gives a mixture of sodium sulphide and barium or
strontium sulphide, and if this mixture be dissolved in water and
the solution evaporated, barium or strontium hydroxide
crystallises out on cooling, and sodium hydrosulphide, NaHS, is
obtained in solution. The hydroxides BaH_{2}O_{2} and SrH_{2}O_{2}
are prepared on a large scale, being applied to many reactions;
for example, strontium hydroxide is prepared for sugar works for
extracting crystallisable sugar from molasses.
We may remark that Boussingault, by igniting barium sulphate in
hydrochloric acid gas, obtained a complete decomposition, with the
formation of barium chloride. Attention should also be turned to
the fact that Grouven, by beating a mixture of charcoal and
strontium sulphate with magnesium and potassium sulphates, showed
the easy decomposability depending on the formation of double
salts, such as SrS,K_{2}S, which are easily soluble in water, and
give a precipitate of strontium carbonate with carbonic anhydride.
In such examples as these we see that the force which binds double
salts may play a part in directing the course of reactions, and
the number of double salts of silica on the earth's surface shows
that nature takes advantage of these forces in her chemical
processes. It is worthy of remark that Buchner (1893), by mixing a
40 per cent. solution of barium acetate with a 60 per cent.
solution of sulphate of alumina, obtained a thick glutinous mass,
which only gave a precipitate of BaSO_{4} after being diluted with
water.
[53] Barium sulphate is sometimes converted into barium chloride in the
following manner: finely-ground barium sulphate is heated with
coal and manganese chloride (the residue from the manufacture of
chlorine). The mass becomes semi-liquid, and when it evolves
carbonic oxide the heating is stopped. The following double
decompositions proceed during this operation: first the carbon
takes up the oxygen from the barium sulphate, and gives sulphide,
BaS, which enters into double decomposition with the chloride of
manganese, MnCl_{2}, forming manganese sulphide, MnS, which is
insoluble in water, and soluble barium chloride. This solution is
easily obtained pure because many foreign impurities, such as
iron, remain in the insoluble portion with the manganese. The
solution of barium chloride is chiefly used for the preparation of
barium sulphate, which is precipitated by sulphuric acid, by which
means _barium sulphate_ is re-formed as a powder. This salt is
characterised by the fact that it is unacted on by the majority of
chemical reagents, is insoluble in water, and is not dissolved by
acids. Owing to this, artificial barium sulphate forms a permanent
white paint which is used instead of (and mixed with) white lead,
and has been termed 'blanc fixé' or 'permanent white.'
The solution of one part of calcium chloride at 20° requires 1·36
part of water, the solution of one part of strontium chloride
requires 1·88 part of water at the same temperature, and the
solution of barium chloride 2·88 parts of water. The solubility of
the bromides and iodides varies in the same proportion. The
chlorides of barium and strontium crystallise out from solution
with great ease in combination with water; they form
BaCl_{2},2H_{2}O and SrCl_{2},6H_{2}O. The latter (which separates
out at 40°) resembles the salts of Ca and Mg in composition, and
Étard (1892) obtained SrCl_{2},2H_{2}O from solutions at 90-130°.
We may also observe that the crystallo-hydrates BaBr_{2},H_{2}O
and BaI_{2},7H_{2}O are known.
[54] The nitrates Sr(NO_{3})_{2} (in the cold its solutions give a
crystallo-hydrate containing 4H_{2}O) and Ba(NO_{3})_{2} are so
very sparingly soluble in water that they separate in considerable
quantities when a solution of sodium nitrate is added to a strong
solution of either barium or strontium chloride. They are obtained
by the action of nitric acid on the carbonates or oxides. 100
parts of water at 15° dissolve 6·5 parts of strontium nitrate and
8·2 parts of barium nitrate, whilst more than 300 parts of calcium
nitrate are soluble at the same temperature. Strontium nitrate
communicates a crimson coloration to the flame of burning
substances, and is therefore frequently used for Bengal fire,
fireworks, and signal lights, for which purpose the salts of
lithium are still better fitted. Calcium nitrate is exceedingly
hygroscopic. Barium nitrate, on the contrary, does not show this
property in the least degree, and in this respect it resembles
potassium nitrate, and is therefore used instead of the latter for
the preparation of a gunpowder which is called 'saxifragin powder'
(76 parts of barium nitrate, 2 parts of nitre, and 22 parts of
charcoal).
[55] The dissociation of the crystallo-hydrate of baryta is given in
Chapter I., Note 65. 100 parts of water dissolve
0° 20° 40° 60° 80°
BaO 1·5 3·5 7·4 18·8 90·8
SrO 0·3 0·7 1·4 3 9
Supersaturated solutions are easily formed.
The anhydrous oxide BaO fuses in the oxyhydrogen flame. When
ignited in the vapour of potassium, the latter takes up the
oxygen; whilst in chlorine, oxygen is separated and barium
chloride formed.
[55 bis] Brugellmann, by heating BaH_{2}O_{2} in a graphite or clay
crucible, obtained BaO in needles, sp. gr. 5·32, and by heating in
a platinum crucible--in crystals belonging to the cubical system,
sp. gr. 5·74. SrO is obtained in the latter form from the nitrate.
The following are the specific gravities of the oxides from
different sources:--
MgO CaO SrO
from RN_{2}O_{6} 3·38 3·25 4·75
" RCO_{3} 3·48 3·26 4·45
" RH_{2}O_{2} 3·41 3·25 4·57
[56] The property of barium oxide of absorbing oxygen when heated, and
giving the peroxide, BaO_{2}, is very characteristic for this
oxide (_see_ Chapter III., Note 7). It only belongs to the
anhydrous oxide. The hydroxide does not absorb oxygen. Peroxides
of calcium and strontium may be obtained by means of hydrogen
peroxide. Barium peroxide is insoluble in water, but is able to
form a hydrate with it, and also to combine with hydrogen
peroxide, forming a very unstable compound having the composition
BaH_{2}O_{4} (obtained by Professor Schöne), which in course of
time evolves oxygen (Chapter IV., Note 21).
Barium oxide is decomposed when heated with potassium; fused barium
chloride is decomposed, as Davy showed, by the action of a galvanic
current, forming metallic _barium_; and Crookes (1862) obtained an
amalgam of barium from which the mercury could easily be driven off,
by heating sodium amalgam in a saturated solution of barium chloride.
Strontium is obtained by the same processes. Both metals are soluble in
mercury, and seem to be non-volatile or only very slightly volatile.
They are both heavier than water; the specific gravity of barium is
3·6, and of strontium 2·5. They both decompose water at the ordinary
temperature, like the metals of the alkalis.
Barium and strontium as saline elements are characterised by their
powerful basic properties, so that they form acid salts with difficulty,
and scarcely form basic salts. On comparing them together and with
calcium, it is evident that the alkaline properties in this group (as
in the group potassium, rubidium, cæsium) increase with the atomic
weight, and this succession clearly shows itself in many of their
corresponding compounds. Thus, for instance, the solubility of the
hydroxides RH_{2}O_{2} and the specific gravity[57] rise in passing from
calcium to strontium and barium, while the solubility of the sulphates
decreases,[58] and therefore in the case of magnesium and beryllium,
as metals whose atomic weights are still less, we should expect the
solubility of the sulphates to be greater, and this is in reality the
case.
[57] Even in solutions a gradual progression in the increase of the
specific gravity shows itself, not only for equivalent solutions
(for instance, RCl_{2} + 200H_{2}O), but even with an equal
percentage composition, as is seen from the curves giving the
specific gravity (water 4° = 10,000) at 15° (for barium chloride,
according to Bourdiakoff's determinations):
BeCl_{2} : S = 9,992 + 67·21_p_ + 0·111_p_^2
CaCl_{2} : S = 9,992 + 80·24_p_ + 0·476_p_^2
SrCl_{2} : S = 9,992 + 85·57_p_ + 0·733_p_^2
BaCl_{2} : S = 9,992 + 86·56_p_ + 0·813_p_^2
[58] One part of calcium sulphate at the ordinary temperature requires
about 500 parts of water for solution, strontium sulphate about
7,000 parts, barium sulphate about 400,000 parts, whilst beryllium
sulphate is easily soluble in water.
Just as in the series of the alkali metals we saw the metals potassium,
rubidium, and cæsium approaching near to each other in their properties,
and allied to them two metals having smaller combining weights--namely,
sodium, and the lightest of all, lithium, which all exhibited certain
peculiar characteristic properties--so also in the case of the metals
of the alkaline earths we find, besides calcium, barium, and strontium,
the metal magnesium and also _beryllium_ or _glucinum_. In respect
to the magnitude of its atomic weight, this last occupies the same
position in the series of the metals of the alkaline earths as lithium
does in the series of the alkali metals, for the combining weight of
beryllium, Be or Gl = 9. This combining weight is greater than that
of lithium (7), as the combining weight of magnesium (24) is greater
than that of sodium (23), and as that of calcium (40) is greater than
that of potassium (39), &c.[59] Beryllium was so named because it
occurs in the mineral _beryl_. The metal is also called glucinum (from
the Greek word [Greek: glykys], 'sweet'), because its salts have a
sweet taste. It occurs in beryl, aquamarine, the emerald, and other
minerals, which are generally of a green colour; they are sometimes
found in considerable masses, but as a rule are comparatively rare
and, as transparent crystals, form precious stones. The composition of
beryl and of the emerald is as follows: Al_{2}O_{3},3BeO,6SiO_{2}. The
Siberian and Brazilian beryls are the best known. The specific gravity
of beryl is about 2·7. Beryllium oxide, from the feebleness of its basic
properties, presents an analogy to aluminium oxide in the same way
that lithium oxide is analogous to magnesium oxide.[60] Owing to its
rare occurrence in nature, to the absence of any especially distinct
individual properties, and to the possibility of foretelling them to
a certain extent on the basis of the periodic system of the elements
given in the following chapter, and owing to the brevity of this
treatise, we will not discuss at any length the compounds of beryllium,
and will only observe that their individuality was pointed out in 1798
by Vauquelin, and that metallic beryllium was obtained by Wöhler and
Bussy. Wöhler obtained _metallic beryllium_ (like magnesium) by acting
on beryllium chloride, BeCl_{2}, with potassium (it is best prepared
by fusing K_{2}BeF_{4} with Na). Metallic beryllium has a specific
gravity 1·64 (Nilson and Pettersson). It is very infusible, melting at
nearly the same temperature as silver, which it resembles in its white
colour and lustre. It is characterised by the fact that it is very
difficultly oxidised, and even in the oxidising flame of the blowpipe is
only superficially covered by a coating of oxide; it does not burn in
pure oxygen, and does not decompose water at the ordinary temperature
or at a red heat, but gaseous hydrochloric acid is decomposed by it
when slightly heated, with evolution of hydrogen and development of a
considerable amount of heat. Even dilute hydrochloric acid acts in the
same manner at the ordinary temperature. Beryllium also acts easily on
sulphuric acid, but it is remarkable that neither dilute nor strong
nitric acid acts on beryllium, which seems especially able to resist
oxidising agents. Potassium hydroxide acts on beryllium as on aluminium,
hydrogen being disengaged and the metal dissolved, but ammonia has no
action on it. These properties of metallic beryllium seem to isolate it
from the series of the other metals described in this chapter, but if we
compare the properties of calcium, magnesium, and beryllium we shall see
that magnesium occupies a position intermediate between the other two.
Whilst calcium decomposes water with great ease, magnesium does so with
difficulty, and beryllium not at all. The peculiarities of beryllium
among the metals of the alkaline earths recall the fact that in the
series of the halogens we saw that fluorine differed from the other
halogens in many of its properties and had the smallest atomic weight.
The same is the case with regard to beryllium among the other metals of
the alkaline earths.
[59] We refer beryllium to the class of the bivalent metals of the
alkaline earths--that is, we ascribe to its oxide the formula BeO,
and do not consider it as trivalent (Be = 13·5, Chapter VII., Note
21), although that view has been upheld by many chemists. The true
atomic composition of beryllium oxide was first given by the
Russian chemist, Avdéeff (1819), in his researches on the
compounds of this metal. He compared the compounds of beryllium to
those of magnesium, and refuted the notion prevalent at the time,
of the resemblance between the oxides of beryllium and aluminium,
by proving that beryllium sulphate presents a greater resemblance
to magnesium sulphate than to aluminium sulphate. It was
especially noticed that the analogues of alumina give alums,
whilst beryllium oxide, although it is a feeble base, easily
giving, like magnesia, basic and double salts, does not form true
alums. The establishment of the periodic system of the elements
(1869), considered in the following chapter, immediately indicated
that Avdéeff's view corresponded with the truth--that is, that
beryllium is bivalent, which therefore necessitated the denial of
its trivalency. This scientific controversy resulted in a long
series of researches (1870-80) concerning this element, and ended
in Nilson and Pettersson--two of the chief advocates of the
trivalency of beryllium--determining the vapour density of
BeCl_{2} = 40, (Chapter VII., Note 21), which gave an undoubted
proof of its bivalency (_see_ also Note 3).
[60] Beryllium oxide, like aluminium oxide, is precipitated from
solutions of its salts by alkalis as a gelatinous hydroxide,
BeH_{2}O_{2}, which, like alumina, is soluble in an excess of
caustic potash or soda. This reaction may be taken advantage of
for distinguishing and separating beryllium from aluminium,
because when the alkaline solution is diluted with water and
boiled, beryllium hydroxide is precipitated, whilst the alumina
remains in solution. The solubility of the beryllium oxide at once
clearly indicates its feeble basic properties, and, as it were,
separates this oxide from the class of the alkaline earths. But on
arranging the oxides of the above-described metals of the alkaline
earths according to their decreasing atomic weights we have the
series
BaO, SrO, CaO, MgO, BeO,
in which the basic properties and solubility of the oxides
consecutively and distinctly decrease until we reach a point when,
had we not known of the existence of the beryllium oxide, we
should expect to find in its place an oxide insoluble in water and
of feeble basic properties. If an alcoholic solution of caustic
potash be saturated with the hydrate of BeO, and evaporated under
the receiver of an air pump, it forms silky crystals BeK_{2}O_{2}.
Another characteristic of the salts of beryllium is that they give
with aqueous ammonia a gelatinous precipitate which is soluble in
an excess of ammonium carbonate like the precipitate of magnesia;
in this beryllium oxide differs from the oxide of aluminium.
Beryllium oxide easily forms a carbonate which is insoluble in
water, and resembles magnesium carbonate in many respects.
Beryllium sulphate is distinguished by its considerable solubility
in water--thus, at the ordinary temperature it dissolves in an
equal weight of water; it crystallises out from its solutions in
well-formed crystals, which do not change in the air, and contain
BeSO_{4},4H_{2}O. When ignited it leaves beryllium oxide, but this
oxide, after prolonged ignition, is re-dissolved by sulphuric
acid, whilst aluminium sulphate, after a similar treatment, leaves
aluminium oxide, which is no longer soluble in acids. With a few
exceptions, the salts of beryllium crystallise with great
difficulty, and to a considerable extent resemble the salts of
magnesium; thus, for instance, beryllium chloride is analogous to
magnesium chloride. It is volatile in an anhydrous state, and in a
hydrated state it decomposes, with the evolution of hydrochloric
acid.
In addition to the above characteristics of the compounds of the metals
of the alkaline earths, we must add that they, like the alkali metals,
combine with nitrogen and hydrogen, and while sodium nitride (obtained
by igniting the amide of sodium, Chapter XII., Note 44 bis) and lithium
nitride (obtained by heating lithium in nitrogen, Chapter XIII., Note
39) have the composition R_{3}N, so the nitrides of magnesium (Note 14),
calcium, strontium, and barium have the composition R_{3}N_{2}, for
example, Ba_{3}N_{2}, as might be expected from the diatomicity of the
metals of the alkaline earths and from the relation of the nitrides to
ammonia, which is obtained from all of these compounds by the action of
water. The _nitrides_ of Ca, Sr, and Ba are formed directly (Maquenne,
1892) by heating the metals in nitrogen. They all have the appearance of
an amorphous powder of dark colour; as regards their reactions, it is
known that besides disengaging ammonia with water, they form cyanides
when heated with carbonic oxide; for instance, Ba_{3}N_{2} + 2CO =
Ba(CN)_{2} + 2BaO.[61]
[61] Thus in the nitrides of the metals we have substances by means
of which we can easily obtain from the nitrogen of the air, not
only ammonia, but also with the aid of CO, by synthesis, a whole
series of complex carbon and nitrogen compounds.
The metals of the alkaline earths, just like Na and K, absorb hydrogen
under certain conditions, and form pulverulent easily oxidisable
metallic hydrides, whose composition corresponds exactly to that of
Na_{2}H and K_{2}H, with the substitution of K_{2} and Na_{2} by the
atoms Be, Mg, Ca, Sr, and Ba. The _hydrides of the metals of the
alkaline earths_ were discovered by C. Winkler (1891) in investigating
the reducibility of these metals by magnesium. In reducing their oxides
by heating them with magnesium powder in a stream of hydrogen, Winkler
observed that the hydrogen was absorbed (but very slowly), _i.e._ at
the moment of their separation all the metals of the alkaline earths
combine with hydrogen. This absorptive power increases in passing from
Be to Mg, Ca, Sr, and Ba, and the resultant hydrides retain the combined
hydrogen[62] when heated, so that these hydrides are distinguished
for their considerable stability under heat, but they oxidise very
easily.[63]
[62] As the hydrides of calcium, magnesium, &c. are very stable under
the action of heat, and these metals and hydrogen occur in the
sun, it is likely that the formation of their hydrides may take
place there. (Private communication from Prof. Winkler, 1894.) It
is probable that in the free metals of the alkaline earths
hitherto obtained a portion was frequently in combination with
nitrogen and hydrogen.