(Chapter VIII., Note 1), contains little potash. For the extraction
of potash, which was formerly carried on extensively in the east of
Russia (before the discovery of the Stassfurt salt), the ash of
grasses, and the green portions of potatoes, buckwheat, &c., are
taken and treated with water (lixiviated), the solution is
evaporated, and the residue ignited in order to destroy the organic
matter present in the extract. The residue thus obtained is
composed of raw potash. It is refined by a second dissolution in a
small quantity of water, for the potash itself is very soluble in
water, whilst the impurities are sparingly soluble. The solution
thus obtained is again evaporated, and the residue ignited, and
this potash is then called refined potash, or pearlash. This method
of treatment cannot give chemically pure potassium carbonate. A
certain amount of impurities remain. To obtain chemically pure
potassium carbonate, some other salt of potassium is generally
taken and purified by crystallisation. Potassium carbonate
crystallises with difficulty, and it cannot therefore be purified
by this means, whilst other salts, such as the tartrate, acid
carbonate, sulphate, or nitrate, &c., crystallise easily and may
thus be directly purified. The tartrate is most frequently
employed, since it is prepared in large quantities (as a sediment
from wine) for medicinal use under the name of cream of tartar.
When ignited without access of air, it leaves a mixture of charcoal
and potassium carbonate. The charcoal so obtained being in a
finely-divided condition, the mixture (called 'black flux'), is
sometimes used for reducing metals from their oxides with the aid
of heat. A certain quantity of nitre is added to burn the charcoal
formed by heating the cream of tartar. Potassium carbonate thus
prepared is further purified by converting it into the acid salt,
by passing a current of carbonic anhydride through a strong
solution. KHCO_{3} is then formed, which is less soluble than the
normal salt (as is also the case with the corresponding sodium
salts), and therefore crystals of the acid salt separate from the
solution on cooling. When ignited, they part with their water and
carbonic anhydride, and pure potassium carbonate remains behind.
The physical properties of potassium carbonate distinguish it
sufficiently from sodium carbonate; it is obtained from solutions
as a powdery white mass, having an alkaline taste and reaction,
and, as a rule, shows only traces of crystallisation. It also
attracts the moisture of the air with great energy. The crystals do
not contain water, but absorb it from the air, deliquescing into a
saturated solution. It melts at a red heat (1045°), and at a still
higher temperature is even converted into vapour, as has been
observed at glass works where it is employed. It is very soluble.
At the ordinary temperature, water dissolves an equal weight of the
salt. Crystals containing two equivalents of water separate from
such a saturated solution when strongly cooled (Morel obtained
K_{2}CO_{3}3H_{2}O in well-formed crystals at +10°). There is no
necessity to describe its reactions, because they are all analogous
to those of sodium carbonate. When manufactured sodium carbonate
was but little known, the consumption of potassium carbonate was
very considerable, and even now washing soda is frequently replaced
for household purposes by 'lye'--_i.e._ an aqueous solution
obtained from ashes. It contains potassium carbonate, which acts
like the sodium salt in washing tissues, linen, &c.
A mixture of potassium and sodium carbonates fuses with much
greater ease than the separate salts, and a mixture of their
solutions gives well-crystallised salts--for instance (Marguerite's
salt), K_{2}CO_{3},6H_{2}O,2Na_{2}CO_{3},6H_{2}O. Crystallisation
also occurs in other multiple proportions of K and Na (in the above
case 1 : 2, but 1 : 1 and 1 : 3 are known), and always with 6 mol.
H_{2}O. This is evidently a combination _by similarity_, as in
alloys, solutions, &c.
[8 bis] About 25,000 tons of potash annually are now prepared from KCl
by this method at Stassfurt.
[9] _Potassium sulphate_, K_{2}SO_{4}, crystallises from its solutions
in an anhydrous condition, in which respect it differs from the
corresponding sodium salt, just as potassium carbonate differs from
sodium carbonate. In general, it must be observed that the majority
of sodium salts combine more easily with water of crystallisation
than the potassium salts. The solubility of _potassium sulphate_
does not show the same peculiarities as that of sodium sulphate,
because it does not combine with water of crystallisation; at the
ordinary temperature 100 parts of water dissolve about 10 parts of
the salt, at 0° 8·3 parts, and at 100° about 26 parts. _The acid
sulphate_, KHSO_{4}, obtained easily by heating crystals of the
normal salt with sulphuric acid, is frequently employed in chemical
practice. On heating the mixture of acid and salt, fumes of
sulphuric acid are at first given off; when they cease to be
evolved, the acid salt is contained in the residue. At a higher
temperature (of above 600°) the acid salt parts with all the acid
contained in it, the normal salt being re-formed. The definite
composition of this acid salt, and the ease with which it
decomposes, render it exceedingly valuable for certain chemical
transformations accomplished by means of sulphuric acid at a high
temperature, because it is possible to take, in the form of this
salt, a strictly definite quantity of sulphuric acid, and to cause
it to act on a given substance at a high temperature, which it is
often necessary to do, more especially in chemical analysis. In
this case, the acid salt acts in exactly the same manner as
sulphuric acid itself, but the latter is inefficient at
temperatures above 400°, because it all evaporates, while at that
temperature the acid salt still remains in a fused state, and acts
with the elements of sulphuric acid on the substance taken. Hence
by its means the boiling-point of sulphuric acid is raised. Thus
the acid potassium sulphate is employed, where for conversion of
certain oxides, such as those of iron, aluminium, and chromium,
into salts, a high temperature is required.
Weber, by heating potassium sulphate with an excess of sulphuric
acid at 100°, observed the formation of a lower stratum, which was
found to contain a definite compound containing eight equivalents
of SO_{3} per equivalent of K_{2}O. The salts of rubidium, cæsium,
and thallium give a similar result, but those of sodium and lithium
do not. (_See_ Note 1.)
[10] The _bromide_ and _iodide_ of potassium are used, like the
corresponding sodium compounds, in medicine and photography.
Potassium iodide is easily obtained in a pure state by saturating
a solution of hydriodic acid with caustic potash. In practice,
however, this method is rarely had recourse to, other more simple
processes being employed although they do not give so pure a
product. They aim at the direct formation of hydriodic acid in the
liquid in the presence of potassium hydroxide or carbonate. Thus
iodine is thrown into a solution of pure potash, and hydrogen
sulphide passed through the mixture, the iodine being thus
converted into hydriodic acid. Or a solution is prepared from
phosphorus, iodine, and water, containing hydriodic and phosphoric
acid; lime is then added to this solution, when calcium iodide is
obtained in solution, and calcium phosphate as a precipitate. The
solution of calcium iodide gives, with potassium carbonate,
insoluble calcium carbonate and a solution of potassium iodide. If
iodine is added to a slightly-heated solution of caustic potash
(free from carbonate--that is, freshly prepared), so long as the
solution is not coloured from the presence of an excess of iodine,
there is formed (as in the action of chlorine on a solution of
caustic potash) a mixture of potassium iodide and iodate. On
evaporating the solution thus obtained and igniting the residue,
the iodate is destroyed and converted into iodide, the oxygen
being disengaged, and potassium iodide only is left behind. On
dissolving the residue in water and then evaporating, cubical
crystals of the anhydrous salt are obtained, which are soluble in
water and alcohol, and on fusion give an alkaline reaction, owing
to the fact that when ignited a portion of the salt decomposes,
forming potassium oxide. The neutral salt may be obtained by
adding hydriodic acid to this alkaline salt until it gives an acid
reaction. It is best to add some finely-divided charcoal to the
mixture of iodate and iodide before igniting it, as this
facilitates the evolution of the oxygen from the iodate. The
iodate may also be converted into iodide by the action of certain
reducing agents, such as zinc amalgam, which when boiled with a
solution containing an iodate converts it into iodide. Potassium
iodide may also be prepared by mixing a solution of ferrous iodide
(it is best if the solution contain an excess of iodine) and
potassium carbonate, in which case ferrous carbonate FeCO_{3}, is
precipitated (with an excess of iodine the precipitate is
granular, and contains a compound of the suboxide and oxide of
iron), while potassium iodide remains in solution. Ferrous iodide,
FeI_{2}, is obtained by the direct action of iodine on iron in
water. Potassium iodide considerably lowers the temperature (by
24°), when it dissolves in water, 100 parts of the salt dissolve
in 73·5 parts of water at 12·5°, in 70 parts at 18°, whilst the
saturated solution which boils at 120° contains 100 parts of salt
per 45 parts of water. Solutions of potassium iodide dissolve a
considerable amount of iodine; strong solutions even dissolving as
much or more iodine than they contain as potassium iodide (_see_
Note 3 bis and Chapter XI., Note 64).
[11] Caustic potash is not only formed by the action of lime on dilute
solutions of potassium carbonate (as sodium hydroxide is prepared
from sodium carbonate), but by igniting potassium nitrate with
finely-divided copper (_see_ Note 15), and also by mixing
solutions of potassium sulphate (or even of alum, KAlS_{2}O_{8})
and barium hydroxide, BaH_{2}O_{2}. It is sometimes purified by
dissolving it in alcohol (the impurities, for example, potassium
sulphate and carbonate, are not dissolved) and then evaporating
the alcohol.
The specific gravity of potassium hydroxide is 2·04, but that of
its solutions (see Chapter XII., Note 18) at 15° S = 9,992 +
90·4_p_ + 0·28_p_^2 (here _p_^2 is +, and for sodium hydroxide it
is-). Strong solutions, when cooled, yield a crystallo-hydrate,
KHO,4H_{2}O, which dissolves in water, producing cold (like
2NaHO,7H_{2}O), whilst potassium hydroxide in solution develops a
considerable amount of heat.
_Potassium cyanide_, which presents in its chemical relations a certain
analogy with the halogen salts of potassium, is not only formed
according to the equation, KHO + HCN = H_{2}O + KCN, but also whenever
a nitrogenous carbon compound--for instance, animal matter--is heated
in the presence of metallic potassium, or of a compound of potassium,
and even when a mixture of potash and carbon is heated in a stream of
nitrogen. Potassium cyanide is obtained from yellow prussiate, which
has been already mentioned in Chapter IX., and whose preparation
on a large scale will be described in Chapter XXII. If the yellow
prussiate be ground to a powder and dried, so that it loses its water
of crystallisation, it then melts at a red heat, and decomposes into
carbide of iron, nitrogen, and potassium cyanide, FeK_{4}C_{6}N_{6} =
4KCN + FeC_{2} + N_{2}. After the decomposition it is found that the
yellow salt has been converted into a white mass of potassium cyanide.
The carbide of iron formed collects at the bottom of the vessel. If
the mass thus obtained be treated with water, the potassium cyanide is
partially decomposed by the water, but if it be treated with alcohol,
then the cyanide is dissolved, and on cooling separates in a crystalline
form.[12] A solution of potassium cyanide has a powerfully alkaline
reaction, a smell like that of bitter almonds, peculiar to prussic acid,
and acts as a most powerful poison. Although exceedingly stable in a
fused state, potassium cyanide easily changes when in solution. Prussic
acid is so very feebly energetic that even water decomposes potassium
cyanide. A solution of the salt, even without access of air, easily
turns brown and decomposes, and when heated evolves ammonia and forms
potassium formate; this is easily comprehensible from the representation
of the cyanogen compounds which was developed in Chapter IX., KCN +
2H_{2}O = CHKO_{2} + NH_{3}. Furthermore, as carbonic anhydride acts on
potassium cyanide with evolution of prussic acid, and as potassium
cyanate, which is also unstable, is formed by the action of air, it will
be easily seen that solutions of potassium cyanide are very unstable.
Potassium cyanide, containing as it does carbon and potassium, is a
substance which can act in a very vigorously reducing manner, especially
when fused; it is therefore used as a powerful reducing agent at a red
heat.[13] The property of potassium cyanide of giving double salts
with other cyanides is very clearly shown by the fact that many metals
dissolve in a solution of potassium cyanide, with the evolution of
hydrogen. For example, iron, copper, and zinc act in this manner. Thus--
4KCN + 2H_{2}O + Zn = K_{2}ZnC_{4}N_{4} + 2KHO + H_{2}
[12] When the yellow prussiate is heated to redness, all the cyanogen
which was in combination with the iron is decomposed into
nitrogen, which is evolved as gas, and carbon, which combines with
the iron. In order to avoid this, potassium carbonate is added to
the yellow prussiate while it is being fused. A mixture of 8 parts
of anhydrous yellow prussiate and 3 parts of pure potassium
carbonate is generally taken. Double decomposition then takes
place, resulting in the formation of ferrous carbonate and
potassium cyanide. But by this method, as by the first, a pure
salt is not obtained, because a portion of the potassium cyanide
is oxidised at the expense of the iron carbonate and forms
potassium cyanate, FeCO_{3} + KCN = CO_{2} + Fe + KCNO; and the
potassium cyanide very easily forms oxide, which acts on the sides
of the vessel in which the mixture is heated (to avoid this iron
vessels should be used). By adding one part of charcoal powder to
the mixture of 8 parts of anhydrous yellow prussiate and 3 parts
of potassium carbonate a mass is obtained which is free from
cyanate, because the carbon absorbs the oxygen, but in that case
it is impossible to obtain a colourless potassium cyanide by
simple fusion, although this may be easily done by dissolving it
in alcohol. Cyanide of potassium may also be obtained from
potassium thiocyanate, which is formed from ammonium thiocyanate
obtained by the action of ammonia upon bisulphide of carbon (_see_
works upon Organic Chemistry). Potassium cyanide is now prepared
in large quantities from yellow prussiate for gilding and
silvering. When fused in large quantities the action of the oxygen
of the air is limited, and with great care the operation may be
successfully conducted, and therefore, on a large scale, very pure
salt is sometimes obtained. When slowly cooled, the fused salt
separates in cubical crystals like potassium chloride.
Pure KCN is obtained by passing CNH gas into an alcoholic solution
of KHO. The large amount of potassium cyanide which is now
required for the extraction of gold from its ores, is being
replaced by a mixture (Rossler and Gasslaker, 1892) of KCN and
NaCN, prepared by heating powdered and dried yellow prussiate with
metallic sodium: K_{4}Fe(CN)_{6} + 2Na = 4KCN + 2NaCN + Fe. This
method offers two advantages over the above methods: (1) the whole
of the cyanide is obtained, and does not decompose with the
formation of N_{2}; and (2) no cyanates are formed, as is the case
when carbonate of potash is heated with the prussiate.
[13] A considerable quantity of potassium cyanide is used in the arts,
more particularly for the preparation of metallic solutions which
are decomposed by the action of a galvanic current; thus it is
very frequently employed in electro-silvering and gilding. An
alkaline solution is prepared, which is moderately stable owing to
the fact that potassium cyanide in the form of certain double
salts--that is, combined with other cyanides--is far more stable
than when alone (yellow prussiate, which contains potassium
cyanide in combination with ferrous cyanide, is an example of
this).
Gold and silver are soluble in potassium cyanide in the presence of air,
in which case the hydrogen, which would otherwise be evolved in the
reaction, combines with the oxygen of the air, forming water (Eissler,
MacLaurin, 1893), for example, 4Au + 4KCN + O + H_{2}O = 2AuKC_{2}N_{2}
+ 2KHO, which is taken advantage of for extracting gold from its ores
(Chapter XXIV.).[13 bis] Platinum, mercury, and tin are not dissolved in
a solution of potassium cyanide, even with access of air.
[13 bis] A dilute solution of KCN is taken, not containing more than
1 per cent. KCN. MacLaurin explains this by the fact that strong
solutions dissolve gold less rapidly, owing to their dissolving
less air, whose oxygen is necessary for the reaction.
_Potassium nitrate_, or common _nitre_ or _saltpetre_, KNO_{3}, is
chiefly used as a component part of gunpowder, in which it cannot be
replaced by the sodium salt, because the latter is deliquescent. It
is necessary that the nitre in gunpowder should be perfectly pure, as
even small traces of sodium, magnesium, and calcium salts, especially
chlorides, render the nitre and the gunpowder capable of attracting
moisture. Nitre may easily be obtained pure, owing to its great
disposition to form crystals both large and small, which aids its
separation from other salts. The considerable differences between the
solubility of nitre at different temperatures aids this crystallisation.
A solution of nitre saturated at its boiling point (116°) contains
335 parts of nitre to 100 parts of water, whilst at the ordinary
temperature--for instance, 20°--the solution is only able to retain 32
parts of the salt. Therefore, in the preparation and refining of nitre,
its solution, saturated at the boiling point, is cooled, and nearly
all the nitre is obtained in the form of crystals. If the solution be
quietly and slowly cooled in large quantities then large crystals are
formed, but if it be rapidly cooled and agitated then small crystals are
obtained. In this manner, if not all, at all events the majority, of
the impurities present in small quantities remain in the mother liquor.
If an unsaturated solution of nitre be rapidly cooled, so as to prevent
the formation of large crystals (in whose crevices the mother liquor,
together with the impurities, would remain), the very minute crystals of
nitre known as saltpetre flour are obtained.
Common nitre occurs in nature, but only in small quantities in admixture
with other nitrates, and especially with sodium, magnesium, and calcium
nitrates. Such a mixture of salts of nitric acid is formed in nature
in fertile earth, and in those localities where, as in _the soil_,
nitrogenous organic remains are decomposed in the presence of alkalis or
alkaline bases with free access of air. This method of the formation of
nitrates requires moisture, besides the free access of air, and takes
place principally during warm weather.[14] In warm countries, and in
temperate climates during the summer months, fertile soils produce a
small quantity of nitre. In this respect India is especially known as
affording a considerable supply of nitre extracted from the soil. The
nitre-bearing soil after the rainy season sometimes becomes covered
during the summer with crystals of nitre, formed by the evaporation of
the water in which it was previously dissolved. This soil is collected,
subjected to repeated lixiviations, and treated for nitre as will be
presently described. In temperate climates nitrates are obtained from
the lime rubbish of demolished buildings which have stood for many
years, and especially from those portions which have been in contact
with the ground. The conditions there are very favourable for the
formation of nitre, because the lime used as a cement in buildings
contains the base necessary for the formation of nitrates, while the
excrement, urine, and animal refuse are sources of nitrogen. By the
methodical lixiviation of this kind of rubbish a solution of nitrogenous
salts is formed similar to that obtained by the lixiviation of fertile
soil. A similar solution is also obtained by the lixiviation of the
so-called _nitre plantations_. They are composed of manure interlaid
with brushwood, and strewn over with ashes, lime, and other alkaline
rubbish. These nitre plantations are set up in those localities where
the manure is not required for the fertilisation of the soil, as, for
example, in the south-eastern 'black earth' Governments of Russia.
The same process of oxidation of nitrogenous matter freely exposed to
air and moisture during the warm season in the presence of alkalis
takes place in nitre plantations as in fertile soil and in the walls
of buildings. From all these sources there is obtained a solution
containing various salts of nitric acid mixed with soluble organic
matter. The simplest method of treating this impure solution of nitre
is to add a solution of potassium carbonate, or to simply treat it with
ashes containing this substance. The potassium carbonate enters into
double decomposition with the calcium and magnesium salts, forming
insoluble carbonates of these bases and leaving the nitre in solution.
Thus, for instance, K_{2}CO_{3} + Ca(NO_{3})_{2} = 2KNO_3 + CaCO_{3}.
Both calcium and magnesium carbonates are insoluble, and therefore after
treatment with potassium carbonate the solution no longer contains salts
of these metals but only the salts of sodium and potassium together with
organic matter. The latter partially separates on heating in an insoluble
form, and is entirely destroyed by heating the nitre to a low red heat.
The nitre thus obtained is easily purified by repeated crystallisation.
The greater part of the nitre used for making gunpowder is now obtained
from the sodium salt _Chili saltpetre_ or _cubic nitre_, which occurs in
nature, as already mentioned. The conversion of this salt into common
nitre is also carried on by means of a double decomposition. This is done
either by adding potassium carbonate (when, on mixing the strong and hot
solutions, sodium carbonate is directly obtained as a precipitate), or,
as is now most frequent, potassium chloride. When a mixture of strong
solutions of potassium chloride and sodium nitrate is evaporated, sodium
chloride first separates, because this salt, which is formed by the
double decomposition KCl + NaNO_{3} = KNO_{3} + NaCl, is almost equally
soluble in hot and cold water; on cooling, therefore, a large amount of
potassium nitrate separates from the saturated solution, while the sodium
chloride remains dissolved. The nitre is ultimately purified by
recrystallisation and by washing with a saturated solution of nitre,
which cannot dissolve a further quantity of nitre but only the
impurities.
[14] Besides which Schloesing and Müntz, by employing similar methods
to Pasteur, showed that the formation of nitre in the
decomposition of nitrogenous substances is accomplished by the aid
of peculiar micro-organisms (ferments), without which the
simultaneous action of the other necessary conditions (alkalis,
moisture, a temperature of 37°, air, and nitrogenous substances)
cannot give nitre.
Nitre is a colourless salt having a peculiar cool taste. It crystallises
easily in long striated six-sided rhombic prisms terminating in rhombic
pyramids. Its crystals (sp. gr. 1·93) do not contain water, but their
cavities generally contain a certain quantity of the solution from
which they have crystallised. For this reason in refining nitre,
the production of large crystals is prevented, _saltpetre flour_
being prepared. At a low red heat (339°) nitre melts to a colourless
liquid.[14 bis] Potassium nitrate at the ordinary temperature and in a
solid form is inactive and stable, but _at a high temperature_ it acts
as a powerful _oxidising agent_, giving up a considerable amount of
oxygen to substances mixed with it.[15] When thrown on to incandescent
charcoal it brings about its rapid combustion, and a mechanical mixture
of powdered charcoal and nitre ignites when brought into contact
with a red-hot substance, and continues to burn by itself. In this
action, nitrogen is evolved, and the oxygen oxidises the charcoal, in
consequence of which potassium carbonate and carbonic anhydride are
formed: 4KNO_{3} + 5C = 2K_{2}CO_{3} + 3CO_{2} + 2N_{2}. This phenomenon
depends on the fact that oxygen in combining with carbon evolves more
heat than it does in combining with nitrogen. Hence, when once the
combustion has been started at the expense of the nitre, it is able to
go on without requiring the aid of external heat. A similar oxidation
or combustion at the expense of the contained oxygen takes place when
nitre is heated with different combustible substances. If a mixture of
sulphur and nitre be thrown upon a red-hot surface, the sulphur burns,
forming potassium sulphate and sulphurous anhydride. In this case, also,
the nitrogen of the nitre is evolved as gas: 2KNO_{3} + 2S = K_{2}SO_{4}
+ N_{2} + SO_{2}. A similar phenomenon occurs when nitre is heated with
many metals. The oxidation of those metals which are able to form acid
oxides with an excess of oxygen is especially remarkable. In this case
they remain in combination with potassium oxide as potassium salts.
Manganese, antimony, arsenic, iron, chromium, &c. are instances of this
kind. These elements, like carbon and sulphur, displace free nitrogen.
The lower oxides of these metals when fused with nitre pass into the
higher oxides. Organic substances are also oxidised when heated with
nitre--that is, they burn at the expense of the nitre. It will be
readily understood from this that nitre is frequently used in practical
chemistry and the arts as an oxidising agent at high temperatures.
Its application in _gunpowder_ is based on this property; gunpowder
consists of a mechanical mixture of finely-ground sulphur, nitre, and
charcoal. The relative proportion of these substances varies according
to the destination of the powder and to the kind of charcoal employed
(a friable, incompletely-burnt charcoal, containing therefore hydrogen
and oxygen, is employed). Gases are formed in its combustion, chiefly
nitrogen and carbonic anhydride, which create a considerable pressure
if their escape be in any way impeded. This action of gunpowder may be
expressed by the equation: 2KNO_{3} + 3C + S = K_{2}S + 3CO_{2} + N_{2}.
[14 bis] Before fusing, the crystals of potassium nitrate change their
form, and take the same form as sodium nitrate--that is, they
change into rhombohedra. Nitre crystallises from hot solutions,
and in general under the influence of a rise of temperature, in a
different form from that given at the ordinary or lower
temperatures. Fused nitre solidifies to a radiated crystalline
mass; but it does not exhibit this structure if metallic chlorides
be present, so that this method may be taken advantage of to
determine the degree of purity of nitre.
Carnelley and Thomson (1888) determined the fusing point of
mixtures of potassium and sodium nitrates. The first salt fuses at
339° and the second at 316°, and if _p_ be the percentage amount
of potassium nitrate, then the results obtained were--
_p_ = 10 20 30 40 50 60 70 80 90
298° 283° 268° 242° 231° 231° 242° 284° 306°
which confirms Shaffgotsch's observation (1857) that the lowest
fusing point (about 231°) is given by mixing molecular quantities
(_p_ = 54·3) of the salts--that is, in the formation of the alloy,
KNO_{3},NaNO_{3}.
A somewhat similar result was discovered by the same observers for
the solubility of mixtures of these salts at 20° in 100 parts of
water. Thus, if _p_ be the weight of potassium nitrate mixed with
100-_p_ parts by weight of sodium nitrate taken for solution, and
_c_ be the quantity of the mixed salts which dissolves in 100, the
solubility of sodium nitrate being 85, and of potassium nitrate
34, parts in 100 parts of water, then--
_p_ = 10 20 30 40 50 60 70 80 90
_c_ = 110 136 136 138 106 81 73 54 41
The maximum solubility proved not to correspond with the most
fusible mixture, but to one much richer in sodium nitrate.
Both these phenomena show that in homogeneous liquid mixtures the
chemical forces that act between substances are the same as those
that determine the molecular weights of substances, even when the
mixture consists of such analogous substances as potassium and
sodium nitrates, between which there is no direct chemical
interchange. It is instructive to note also that the maximum
solubility does not correspond with the minimum fusing point,
which naturally depends on the fact that in solution a third
substance, namely water, plays a part, although an attraction
between the salts, like that which exists between sodium and
potassium carbonates (Note 8), also partially acts.
[15] Fused nitre, with a further rise of temperature, disengages oxygen
and then nitrogen. The nitrite KNO_{2} is first formed and then
potassium oxide. The admixture of certain metals--for example, of
finely-divided copper--aids the last decomposition. The oxygen in
this case naturally passes over to the metal.
It is found by this equation that gunpowder should contain thirty-six
parts of charcoal (13·3 p.c.), and thirty-two parts (11·9 p.c.) of
sulphur, to 202 parts (74·8 p.c.) of nitre, which is very near to its
actual composition.[16]
[16] In China, where the manufacture of gunpowder has long been carried
on, 75·7 parts of nitre, 14·4 of charcoal, and 9·9 of sulphur are
used. Ordinary powder for sporting purposes contains 80 parts of
nitre, 12 of charcoal, and 8 of sulphur, whilst the gunpowder used
in heavy ordnance contains 75 of nitre, 15 of charcoal, and 10 of
sulphur. Gunpowder explodes when heated to 300°, when struck, or
by contact with a spark. A compact or finely-divided mass of
gunpowder burns slowly and has but little disruptive action,
because it burns gradually. To act properly the gunpowder must
have a definite rate of combustion, so that the pressure should
increase during the passage of the projectile along the barrel of
the fire-arm. This is done by making the powder in large granules
or in the shape of six-sided prisms with holes through them
(prismatic powder).
The products of combustion are of two kinds: (1) gases which
produce the pressure and are the cause of the dynamical action of
gunpowder, and (2) a solid residue, usually of a black colour
owing to its containing unburnt particles of charcoal. Besides
charcoal, the residue generally contains potassium sulphide,
K_{2}S, and a whole series of other salts--for instance, carbonate
and sulphate. It is apparent from this that the combustion of
gunpowder is not so simple as it appears to be from the above
formula, and hence the weight of the residue is also greater than
indicated by that formula. According to the formula, 270 parts of
gunpowder give 110 parts of residue--that is, 100 parts of powder
give 37·4 parts of residue, K_{2}S, whilst in reality the weight
of the residue varies from 40 p.c. to 70 p.c. (generally 52 p.c.).
This difference depends on the fact that so much oxygen (of the
nitre) remains in the residue, and it is evident that if the
residue varies the composition of the gases evolved by the powder
will vary also, and therefore the entire process will be different
in different cases. The difference in the composition of the gases
and residue depends, as the researches of Gay-Lussac, Shishkoff
and Bunsen, Nobel and Abel, Federoff, Debus, &c., show, on the
conditions under which the combustion of the powder proceeds. When
gunpowder burns in an open space, the gaseous products which are
formed do not remain in contact with the residue, and then a
considerable portion of the charcoal entering into the composition
of the powder remains unburnt, because the charcoal burns after
the sulphur at the expense of the oxygen of the nitre. In this
extreme case the commencement of the combustion of the gunpowder
may be expressed by the equation, 2KNO_{3} + 3C + S = 2C +
K_{2}SO_{4} + CO_{2} + N_{2}. The residue in a blank cartridge
often consists of a mixture of C, K_{2}SO_{4}, K_{2}CO_{3}, and
K_{2}S_{2}O_{3}. If the combustion of the gunpowder be impeded--if
it take place in a cartridge in the barrel of a gun--the quantity
of potassium sulphate will first be diminished, then the amount of
sulphite, whilst the amount of carbonic anhydride in the gases and
the amount of potassium sulphide in the residue will increase. The
quantity of charcoal entering into the action will then be also
increased, and hence the amount in the residue will decrease.
Under these circumstances the weight of the residue will be
less--for example, 4K_{2}CO_{3} + 4S = K_{2}SO_{4} + 3K_{2}S +
4CO_{2}. Besides which, carbonic oxide has been found in the
gases, and potassium bisulphide, K_{2}S_{2}, in the residue of
gunpowder. The amount of potassium sulphide, K_{2}S, increases
with the completeness of the combustion, and is formed in the
residue at the expense of the potassium sulphite. In recent times
the knowledge of the action of gunpowder and other explosives has
made much progress, and has developed into a vast province of
artillery science, which, guided by the discoveries of chemistry,
has worked out a 'smokeless powder' which burns without leaving a
residue, and does not therefore give any 'powder smoke' (to hinder
the rapidity of firing and aiming), and at the same time
disengages a greater volume of gas and consequently gives (under
proper conditions of combustion) the possibility of communicating
to the charge a greater initial velocity, and therefore greater
distance, force, and accuracy of aim. Such 'smokeless powder' is
prepared either from the varieties of nitrocellulose (Chapter VI.,
Note 37) or from a mixture of them with nitro-glycerine (_ibid_).
In burning they give, besides steam and nitrogen, generally a
large amount of oxide of carbon (this is a very serious drawback
in all the present forms of smokeless powder, because carbonic
oxide is poisonous), and also CO_{2}, H_{2}, &c.
_Metallic potassium_ was obtained like sodium; first by the action of
a galvanic current, then by reduction of the hydroxide by means of
metallic iron, and lastly, by the action of charcoal on the carbonate
at a high temperature. The behaviour of metallic potassium differs,
however, from that of sodium, because it easily combines with carbonic
oxide, forming an explosive and inflammable mass.[17]
[17] The substances obtained in this case are mentioned in Chapter IX.,
Note 31.
Potassium is quite as volatile as sodium, if not more so. At the
ordinary temperature potassium is even softer than sodium; its
freshly-cut surfaces present a whiter colour than sodium, but, like
the latter, and with even greater ease, it oxidises in moist air. It
is brittle at low temperatures, but is quite soft at 25°, and melts at
58°. At a low red heat (667°, Perkin) it distils without change, forming
a green vapour, whose density,[18] according to A. Scott (1887), is
equal to 19 (if that of hydrogen = 1). This shows that the molecule
of potassium (like that of sodium, mercury, and zinc) contains but one
atom. This is also the case with many other metals, judging by recent
researches.[19] The specific gravity of potassium at 15° is 0·87, and
is therefore less than that of sodium, as is also the case with all
its compounds.[20] Potassium decomposes water with great ease at the
ordinary temperature, evolving 45,000 heat units per atomic weight in
grams. The heat evolved is sufficient to inflame the hydrogen, the flame
being coloured violet from the presence of particles of potassium.[21]
[18] A. Scott (1887) determined the vapour densities of many of the
alkali elements and their compounds in a platinum vessel heated in
a furnace and previously filled with nitrogen. But these, the
first data concerning a subject of great importance, have not yet
been sufficiently fully described, nor have they received as much
attention as could be desired. Taking the density of hydrogen as
unity, Scott found the vapour densities of the following
substances to be--
Na 12·75 (11·5). KI 92 (84).
K 19 (19·5). RbCl 70 (60).
CsCl 89·5 (84·2). CsI 133 (130).
FeCl_{3} 68. AgCl 80 (71·7).
In brackets are given the densities corresponding with the
formulæ, according to Avogadro-Gerhardt's law. This figure is not
given for FeCl_{3}, because in all probability under these
conditions (the temperature at which it was determined) a portion
of the FeCl_{3} was decomposed. If it was not decomposed, then a
density 81 would correspond with the formula FeCl_{3}, and if the
decomposition were Fe_{2}Cl_{6} = 2FeCl_{2} + Cl_{2}, then the
density should be 54. With regard to the silver chloride, there is
reason to think that the platinum decomposed this salt. The
majority of Scott's results so closely correspond with the formulæ
that a better concord cannot be expected in such determinations.
V. Meyer (1887) gives 93 as the density of KI.
[19] The molecules of non-metals are more complex--for instance, H_{2},
O_{3}, Cl_{2}, &c. But arsenic, whose superficial appearance
recalls that of metals, but whose chemical properties approach
more nearly to the non-metals, has a complex molecule containing
As_{4}.
[20] As the atomic weight of potassium is greater than that of sodium,
the volumes of the molecules, or the quotients of the molecular
weight by the specific gravity, for potassium compounds are
greater than those of sodium compounds, because both the
denominator and numerator of the fraction increase. We cite for
comparison the volumes of the corresponding compounds--
Na 24 NaHO 18 NaCl 28 NaNO_{3} 37 Na_{2}SO_{4} 54
K 45 KHO 27 KCl 39 KNO_{3} 48 K_{2}SO_{4} 66
[21] The same precautions must be taken in decomposing water by
potassium as have to be observed with sodium (Chapter II., Note
8).
It must be observed that potassium decomposes carbonic anhydride
and carbonic oxide when heated, the carbon being liberated and the
oxygen taken up by the metal, whilst on the other hand charcoal
takes up oxygen from potassium, as is seen from the preparation of
potassium by heating potash with charcoal, hence the reaction
K_{2}O + C = K_{2} + CO is reversible and the relation is the same
in this case as between hydrogen and zinc.
With regard to the relation of potassium to hydrogen and oxygen, it
is closely analogous to sodium in this respect. Thus, with hydrogen
it forms potassium hydride, K_{2}H (between 200° and 411°), and with
oxygen it gives a suboxide K_{4}O, oxide K_{2}O, and peroxide, only
more oxygen enters into the composition of the latter than in sodium
peroxide; potassium peroxide contains KO_{2}, but it is probable that
in the combustion of potassium an oxide KO is also formed. Potassium,
like sodium, is soluble in mercury.[22] In a word, the relation between
sodium and potassium is as close as that between chlorine and bromine,
or, better still, between fluorine and chlorine, as the atomic weight
of sodium, 23, is as much greater than that of fluorine, 19, as that of
potassium, 39, is greater than that of chlorine, 35·5.
[22] _Potassium_ forms _alloys with sodium_ in all proportions. The
alloys containing 1 and 3 equivalents of potassium to one
equivalent of sodium are _liquids_, like mercury at the ordinary
temperature. Joannis, by determining the amount of heat developed
by these alloys in decomposing water, found the evolution for
Na_{2}K, NaK, NaK_{2} and NaK_{3} to be 44·5, 44·1, 43·8 and 44·4
thousand heat units respectively (for Na 42·6 and for K 45·4). The
formation of the alloy NaK_{2} is therefore accompanied by the
development of heat, whilst the other alloys may be regarded as
solutions of potassium or sodium in this alloy. In any case a fall
of the temperature of fusion is evident in this instance as in the
alloys of nitre (Note 14). The liquid alloy NaK_{2} is now used
for filling thermometers employed for temperatures above 360°,
when mercury boils.
The resemblance between _potassium_ and _sodium_ is so great that _their
compounds_ can only be easily _distinguished_ in the form of certain of
their salts. For instance, the acid potassium tartrate, C_{4}H_{5}KO_{6}
(cream of tartar), is distinguished by its sparing solubility in
water and in alcohol, and in a solution of tartaric acid, whilst the
corresponding sodium salt is easily soluble. Therefore, if a solution of
tartaric acid be added in considerable excess to the solutions of the
majority of potassium salts, a precipitate of the sparingly-soluble acid
salt is formed, which does not occur with salts of sodium. The chlorides
KCl and NaCl in solutions easily give double salts K_{2}PtCl_{6} and
Na_{2}PtCl_{6}, with platinic chloride, PtCl_{4}, and the solubility
of these salts is very different, especially in a mixture of alcohol
and ether. The sodium salt is easily soluble, whilst the potassium salt
is insoluble or almost so, and therefore the reaction with platinic
chloride is that most often used for the separation of potassium from
sodium, as is more fully described in works on analytical chemistry.
It is possible to discover the least traces of these metals in
admixture together, by means of their property of imparting different
colours to _a flame_. The presence of a salt of sodium in a flame is
recognised by a brilliant yellow coloration, and a pure potassium
salt colours a colourless flame violet. However, in the presence of
a sodium salt the pale violet coloration given by a potassium salt
is quite undistinguishable, and it is at first sight impossible in
this case to discover the potassium salt in the presence of that of
sodium. But by decomposing the light given by a flame coloured by these
metals or a mixture of them, by means of a prism, they are both easily
distinguishable, because the yellow light emitted by the sodium salt
depends on a group of light rays having a definite index of refraction
which corresponds with the yellow portion of the solar spectrum, having
the index of refraction of the Fraunhofer line (strictly speaking,
group of lines) D, whilst the salts of potassium give a light from
which these rays are entirely absent, but which contain rays of a red
and violet colour. Therefore, if a potassium salt occur in a flame, on
decomposing the light (after passing it through a narrow slit) by means
of a prism, there will be seen red and violet bands of light situated
at a considerable distance from each other; whilst if a sodium salt be
present a yellow line will also appear. If both metals simultaneously
occur in a flame and emit light, the spectrum lines corresponding to the
potassium and the sodium will appear simultaneously.
[Illustration: FIG. 72.--Spectroscope. The prism and table are covered
with an opaque cover. The spectrum obtained from the flame coloured
by a substance introduced on the wire is viewed through B. A light is
placed before the scale D in order to illuminate the image of the scale
reflected through B by the side of the prism.]
For convenience in carrying on this kind of testing, _spectroscopes_
(fig. 72) are constructed,[23] consisting of a refracting prism and
three tubes placed in the plane of the refracting angle of the prism.
One of the tubes, C, has a vertical slit at the end, giving access to
the light to be tested, which then passes into the tube (collimator),
containing a lens which gives the rays a parallel direction. The rays
of light having passed through the slit, and having become parallel,
are refracted and dispersed in the prism, and the spectrum formed is
observed through the eye-piece of the other telescope B. The third
tube D contains a horizontal transparent scale (at the outer end)
which is divided into equal divisions. The light from a source such
as a gas burner or candle placed before this tube, passes through the
scale, and is reflected on that face of the prism which stands before
the telescope B, so that the image of the scale is seen through this
telescope simultaneously with the spectrum given by the rays passing
through the slit of the tube C. In this manner the image of the scale
and the spectrum given by the source of light under investigation are
seen simultaneously. If the sun's rays be directed through the slit
of the tube C, then the observer looking through the eye-piece of B
will see the solar spectrum, and (if the aperture of the slit be narrow
and the apparatus correctly adjusted) the dark Fraunhofer lines in
it.[24] Small-sized spectroscopes are usually so adjusted that (looking
through B) the violet portion of the spectrum is seen to the right and
the red portion to the left, and the Fraunhofer line D (in the bright
yellow portion of the spectrum) is situated on the 50th division of the
scale.[25] If the light emitted by an incandescent solid--for example,
the Drummond light--be passed through the spectroscope, then all the
colours of the solar spectrum are seen, but not the Fraunhofer lines. To
observe the result given by a flame coloured by various salts a Bunsen
gas burner (or the pale flame of hydrogen gas issuing from a platinum
orifice) giving so pale a flame that its spectrum will be practically
invisible is placed before the slit. If any compound of sodium be placed
in the flame of the gas burner (for which purpose a platinum wire on
whose end sodium chloride is fused is fixed to the stand), then the
flame is coloured yellow, and on looking through the spectroscope the
observer will see a bright _yellow_ line falling upon the 50th division
of the scale, which is seen together with the spectrum in the telescope.
No yellow lines of other refractive index, nor any rays of any other
colour, will be seen, and, therefore, the spectrum corresponding with
sodium compounds consists of yellow rays of that index of refraction
which belong to the Fraunhofer (black) line D of the solar spectrum.
If a potassium salt be introduced into the flame instead of a sodium
salt, then two bands will be seen which are much feebler than the
bright sodium band--namely, one red line near the Fraunhofer line A and
another violet line. Besides which, a pale, almost continuous, spectrum
will be observed in the central portions of the scale. If a mixture
of sodium and potassium salts be now introduced into the flame, three
lines will be seen simultaneously--namely, the red and pale violet
lines of potassium and the yellow line of sodium. In this manner it is
possible, by the aid of the spectroscope, to determine the relation
between the spectra of metals and known portions of the solar spectrum.
The continuity of the latter is interrupted by dark lines (that is,
by an absence of light of a definite index of refraction), termed the
Fraunhofer lines of the solar spectrum. It has been shown by careful
observations (by Fraunhofer, Brewster, Foucault, Ångstrom, Kirchhoff,
Cornu, Lockyer, Dewar, and others) that there exists an exact _agreement
between the spectra_ of certain _metals_ and certain of the _Fraunhofer_
lines. Thus the bright yellow sodium line exactly corresponds with the
dark Fraunhofer line D of the solar spectrum. A similar agreement is
observed in the case of many other metals. This is not an approximate or
chance correlation. In fact, if a spectroscope having a large number of
refracting prisms and a high magnifying power be used, it is seen that
the dark line D of the solar spectrum consists of an entire system of
closely adjacent but definitely situated fine and wide (sharp, distinct)
dark lines,[26] and an exactly similar group of bright lines is obtained
when the yellow sodium line is examined through the same apparatus, so
that each bright sodium line exactly corresponds with a dark line in the
solar spectrum.[26 bis] This conformity of the bright lines formed by
sodium with the dark lines of the solar spectrum cannot be accidental.
This conclusion is further confirmed by the fact that the bright lines
of other metals correspond with dark lines of the solar spectrum. Thus,
for example, a series of sparks passing between the iron electrodes
of a Ruhmkorff coil gives 450 very distinct lines characterising this
metal. All these 450 bright lines, constituting the whole spectrum
corresponding with iron, are repeated, as Kirchhoff showed, in the
solar spectrum as dark Fraunhofer lines which occur in exactly the
same situations as the bright lines in the iron spectrum, just as the
sodium lines correspond with the band D in the solar spectrum. Many
observers have in this manner studied the solar spectrum and the spectra
of different metals simultaneously, and discovered in the former lines
which correspond not only with sodium and iron, but also with many
other metals.[27] The spectra of such elements as hydrogen, oxygen,
nitrogen, and other gases may be observed in the so-called Geissler's
tubes--that is, in glass tubes containing rarefied gases, through which
the discharge of a Ruhmkorff's coil is passed. Thus hydrogen gives a
spectrum composed of three lines--a red line corresponding with the
Fraunhofer line C, a green line corresponding with the line F, and a
violet line corresponding with one of the lines between G and H. Of
these rays the red is the brightest, and therefore the general colour of
luminous hydrogen (with an electric discharge through a Geissler tube)
is reddish.
[23] For accurate measurements and comparative researches more
complicated spectroscopes are required which give a greater
dispersion, and are furnished for this purpose with several
prisms--for example, in Browning's spectroscope the light passes
through six prisms, and then, having undergone an internal total
reflection, passes through the upper portion of the same six
prisms, and again by an internal total reflection passes into the
ocular tube. With such a powerful dispersion the relative position
of the spectral lines may be determined with accuracy. For the
absolute and exact determination of the wave lengths it is
particularly important that the spectroscope should be furnished
with diffraction gratings. The construction of spectroscopes
destined for special purposes (for example, for investigating the
light of stars, or for determining the absorption spectra in
microscopic preparations, &c.) is exceedingly varied. Details of
the subject must be looked for in works on physics and on spectrum
analysis. Among the latter the best known for their completeness
and merit are those of Roscoe, Kayser, Vogel, and Lecoq de
Boisbaudran.
[24] The arrangement of all the parts of the apparatus so as to give
the clearest possible vision and accuracy of observation must
evidently precede every kind of spectroscopic determination.
Details concerning the practical use of the spectroscope must be
looked for in special works on the subject. In this treatise the
reader is supposed to have a certain knowledge of the physical
data respecting the refraction of light, and its dispersion and
diffraction, and the theory of light, which allows of the
determination of the length of the waves of light in absolute
measure on the basis of observations with diffraction gratings,
the distance between whose divisions may be easily measured in
fractions of a millimetre; by such means it is possible to
determine the wave-length of any given ray of light.
[25] In order to give an idea of the size of the scale, we may observe
that the ordinary spectrum extends from the zero of the scale
(where the red portion is situated) to the 170th division (where
the end of the visible violet portion of the spectrum is
situated), and that the Fraunhofer line A (the extreme prominent
line in the red) corresponds with the 17th division of the scale;
the Fraunhofer line F (at the beginning of the blue, near the
green colour) is situated on the 90th division, and the line G,
which is clearly seen in the beginning of the violet portion of
the spectrum, corresponds with the 127th division of the scale.
[26] The two most distinct lines of D, or of sodium, have wave-lengths
of 589·5 and 588·9 millionths of a millimeter, besides which
fainter and fainter lines are seen whose wave-lengths in
millionths of a millimeter are 588·7 and 588·1, 616·0 and 615·4,
515·5 and 515·2, 498·3 and 498·2, &c., according to Liveing and
Dewar.
[26 bis] In the ordinary spectroscopes which are usually employed in
chemical research, one yellow band, which does not split up into
thinner lines, is seen instead of the system of sodium lines,
owing to the small dispersive power of the prism and the width of
the slit of the object tube.
[27] The most accurate investigations made in this respect are carried
on with spectra obtained by diffraction, because in this case the
position of the dark and bright lines does not depend on the index
of refraction of the material of the prism, nor on the dispersive
power of the apparatus. The best--that is, the most general and
accurate--method of expressing the results of such determinations
consists in determining the lengths of the waves corresponding to
the rays of a definite index of refraction. (Sometimes instead of
this the fraction of 1 divided by the square of the wave-length is
given.) We will express this _wave-length_ in _millionth parts of
a millimetre_ (the ten-millionth parts are already doubtful, and
fall within the limits of error). In order to illustrate the
relation between the wave-lengths and the positions of the lines
of the spectrum, we will cite the wave-lengths corresponding with
the chief Fraunhofer lines and colours of the spectrum.
Fraunhofer line A B C D E b
Wave-length 761·0 687·5 656·6 589·5-588·9 527·3 518·7
+-----------------+ +---------+ +----+ +---+
Colour red orange yellow green
Fraunhofer line F G H
Wave-length 486·5 431·0 397·2
+----------+ +----+
Colour blue violet
In the following table are given the _wave-lengths_ of the light
rays (the longest and most distinct, _see_ later) for certain
elements, those in black type being the most clearly defined and
distinct lines, which are easily obtained either in the flame of a
Bunsen's burner, or in Geissler's tubes, or in general, by an
electric discharge. These lines refer to the elements (the lines
of compounds are different, as will be afterwards explained, but
many compounds are decomposed by the flame or by an electric
discharge), and moreover to the elements in an incandescent and
rarefied gaseous state, for the spectra sometimes vary
considerably with a variation of temperature and pressure.
It may be mentioned that the _red_ colour corresponds with lines
having a wave-length of from 780 (with a greater wave-length the
lines are hardly visible, and are ultra red) to 650, the _orange_
from 650 to 590, the _yellow_ from 590 to 520, the _green_ from
520 to 490, the _blue_ from 490 to 420, and the _violet_ from 420
to 380 millionth parts of a millimetre. Beyond 380 the lines are
scarcely visible, and belong to the ultra-violet. For fluorine
Moissan found as many as 13 bright lines from 744 to 623.
In the table (p. 565) which is arranged in conformity with the
image of the spectrum as it is seen (the red lines on the
left-hand and the violet on the right-hand side), the figures in
black type correspond with lines which are so bright and
distinctly visible that they may easily be made use of, both in
determining the relation between the divisions of the scale and
the wave-lengths, and in determining the admixture of a given
element with another. Brackets join those lines between which
several other lines are clearly visible if the dispersive power of
the spectroscope permits distinguishing the neighbouring lines. In
the ordinary laboratory spectroscopes with one prism, even with
all possible precision of arrangement and with a brilliancy of
light permitting the observations being made with a very narrow
aperture, the lines whose wave-lengths only differ by 2-3
millionths of a millimetre, are blurred together; and with a wide
aperture a series of lines differing by even as much as 20
millionths of a millimetre appear as one wide line. With a faint
light (that is, with a small quantity of light entering into the
spectroscope) only the most _brilliant_ lines are clearly visible.
The _length_ of the lines does not always correspond with their
brilliancy. According to Lockyer this length is determined by
placing the carbon electrodes (between which the incandescent
vapours of the metals are formed), not horizontally to the slit
(as they are generally placed, to give more light), but vertically
to it. Then certain lines appear long and others short. As a rule
(Lockyer, Dewar, Cornu), the longest lines are those with which it
is easiest to obtain _reversed_ spectra (_see_ later).
Consequently, these lines are the most characteristic. Only the
longest and most brilliant are given in our table, which is
composed on the basis of a collection of the data at our disposal
for _bright_ spectra of the _incandescent and rarefied vapours of
the elements_. As the spectra change with great variations of
temperature and vapour density (the faint lines become brilliant
whilst the bright lines sometimes disappear), which is
particularly clear from Ciamician's researches on the halogens,
until the method of observation and the theory of the subject are
enlarged, particular theoretical importance should not be given to
the wave-lengths showing the maximum brilliancy, which only
possess a practical significance in the common methods of
spectroscopic observations. In general the spectra of metals are
simpler than those of the halogens, and the latter are variable;
at an increased pressure all spectral lines become broader.
+-------+-------+--------+--------+-------+-------+------+------+
| N_{2} | O_{2} | Cl_{2} | Br_{2} | I_{2} | Pb | Sn | Tl |
+-------+-------+--------+--------+-------+-------+------+------+
| -- | -- | -- | -- | -- | -- | -- | -- |
| -- | -- | -- | -- | -- | -- | -- | -- |
| 662 | -- | -- | -- | -- | -- | -- | -- |
| 632} | -- | -- | 636 | -- | -- | 645 | -- |
| 620} | 615 | -- | -- | 621 | -- | -- | -- |
| 585} | -- | -- | -- | 613 | 605·7 | 580} | -- |
| 574} | -- | 546} | -- | 579} | 560·7 | 556} | -- |
| 544} | 543 | 539} | 544} | 560} | 554·7 | -- | 549 |
| 535} | 533 | 528} | 523} | 545} | 537 | -- | 535 |
| 527} | 516 | 519} | 517 | 506} | -- | -- | -- |
| 516} | 495} | 494} | 479 | -- | -- | -- | 489 |
| -- | 494} | 480} | 470} | -- | -- | -- | -- |
| 457 | 470} | -- | 462} | -- | -- | 452 | -- |
| 442 | 465} | -- | 454 | 445 | -- | -- | -- |
| 436} | 447} | 436} | 437 | -- | -- | -- | -- |
| 426} | 432} | 431} | -- | -- | -- | -- | -- |
| -- | -- | -- | -- | 421 | -- | -- | -- |
| -- | -- | -- | -- | -- | -- | -- | -- |
| 409 | -- | -- | -- | -- | 406 | -- | -- |
| -- | -- | -- | -- | -- | -- | -- | -- |
+-------+-------+--------+--------+-------+-------+------+------+
+-----+-----+-----+-------+-----+------+------+-----+-------+-----+
| In | Ga | Al | Ba | Sr | Ca | Mg | Zn | Cd | Hg |
|-----+-----+-----+-------+-----|------+------|-----+-------+-----+
| -- | -- | -- | -- | -- | -- | -- | -- | -- | -- |
| -- | -- | -- | -- | -- | -- | -- | -- | -- | -- |
| -- | -- | -- | -- | -- | 646 | -- | -- | -- | -- |
| -- | -- | 624 | 649·7 | 641 | 644 | -- | 636 | 643·8 | -- |
| 619 | -- | 623 | 614 | 606 | -- | -- | -- | -- | 615 |
| -- | -- | -- | -- | -- | 612 | -- | -- | -- | 579 |
| -- | -- | 572 | 553·5 | -- | -- | -- | -- | -- | 577 |
| -- | -- | 570 | 549 | 548 | 559 | -- | -- | 537·7 | 546 |
| 525 | -- | -- | -- | 524 | -- | 518 | -- | 533·6 | -- |
| -- | -- | -- | -- | -- | -- | 516 | 492 | 508·5 | -- |
| -- | -- | -- | 493·3 | -- | -- | -- | 481 | -- | -- |
| -- | -- | -- | -- | -- | -- | 471 | 472 | 479·9 | -- |
| 451 | -- | 466 | 455 | 460 | -- | -- | 468 | 467·7 | -- |
| -- | -- | -- | -- | -- | 445} | 448 | -- | -- | -- |
| -- | -- | -- | -- | -- | 442} | -- | -- | -- | 436 |
| -- | -- | -- | -- | 430 | -- | -- | -- | -- | -- |
| -- | 417 | -- | -- | 421 | 423 | -- | -- | -- | -- |
| -- | -- | -- | 413 | -- | -- | -- | -- | -- | -- |
| 410 | 403 | 396 | -- | 408 | 397 | 384} | -- | -- | 404 |
| -- | -- | 394 | -- | -- | 393 | 383} | - -| -- | -- |
+-----+-----+-----+-------+-----+------+------+-----+-------+-----+
+------+------+-----+-------+-----+-------+------+-------+-------+
| Mn | Fe | Cu | Ag | Cs | Rb | K | Na | Li |
|------+------+-----+-------+-----+-------+------+-------+-------+
| -- | -- | -- | -- | -- | 780 | 770 | -- | -- |
| -- | -- | -- | -- | -- | -- | 766 | -- | -- |
| -- | -- | -- | -- | -- | -- | -- | -- | 670·6 |
| -- | 640 | -- | -- | -- | -- | -- | -- | -- |
| 602} | -- | -- | -- | 622 | 629·6 | -- | -- | 610 |
| 601} | 561} | 578 | -- | 600 | -- | 583} | 589·5 | -- |
| 551} | 544} | 570 | 546·4 | -- | -- | 578} | 588·9 | -- |
| 534} | 537 | 522 | -- | -- | -- | -- | -- | -- |
| -- | 532} | 515 | 520·8 | -- | -- | 535 | -- | -- |
| -- | 521} | 511 | -- | -- | -- | 532} | -- | -- |
| 482} | 496} | -- | -- | -- | -- | -- | -- | 497 |
| 471} | 489} | -- | -- | -- | -- | -- | -- | -- |
| -- | -- | -- | -- | 459 | -- | -- | -- | 460·3 |
| -- | 441} | -- | -- | 456 | -- | -- | -- | -- |
| -- | 430} | -- | -- | -- | -- | -- | -- | -- |
| -- | 427 | -- | -- | -- | -- | -- | -- | -- |
| 424} | -- | -- | 421 | -- | 420 | -- | -- | -- |
| 403} | 407} | -- | -- | -- | -- | -- | -- | -- |
| -- | 404} | -- | -- | -- | -- | 404 | -- | -- |
| -- | -- | -- | -- | -- | -- | -- | -- | -- |
+------+------+-----+-------+-----+-------+------+-------+-------+
+-------+
| H_{2} |
+-------+
| -- |
| -- |
| -- |
| 656·2 |
| -- |
| -- |
| -- |
| -- |
| -- |
| -- |
| 486·1 |
| -- |
| -- |
| -- |
| 434 |
| -- |
| -- |
| -- |
| -- |
| -- |
+-------+
[Illustration: FIG. 73.--Absorption spectrum (Lecoq de Boisbaudran) of
salts of didymium in concentrated and dilute solutions.]
The correlation of the Fraunhofer lines with the spectra of metals
depends on the phenomenon of the so-called _reversal of the spectrum_.
This phenomenon consists in this, that instead of the bright spectrum
corresponding with a metal, under certain circumstances a similar dark
spectrum in the form of Fraunhofer lines may be obtained, as will be
explained directly. In order to clearly understand the phenomenon of
reversed spectra, it must be known that when light passes through
certain transparent substances these substances retain rays of a certain
refrangibility. The colour of solutions is a proof of this. Light
which has passed through a yellow solution of a uranium salt contains
no violet rays, and after having passed through a red solution of a
permanganate, does not contain many rays in the yellow, blue, and green
portions of the spectrum. Solutions of copper salts absorb nearly all
red rays. Sometimes colourless solutions also absorb rays of certain
definite refractive indexes, and give _absorption spectra_. Thus
solutions of salts of didymium absorb rays of a certain refrangibility,
and therefore an impression of black lines is received,[28] as shown in
fig. 73. Many vapours (iodine) and gases (nitric peroxide) give similar
spectra. Light which has passed through a deep layer of aqueous vapour,
oxygen, or nitrogen also gives an absorption spectrum. For this reason
the peculiar (winter) dark lines discovered by Brewster are observed in
sunlight, especially in the evening and morning, when the sun's rays
pass through the atmosphere (containing these substances) by a longer
path than at mid-day. It is evident that the Fraunhofer lines may be
ascribed to the absorption of certain rays of light in its passage from
the luminous mass of the sun to the earth. The remarkable progress made
in all spectroscopic research dates from the investigations made by
_Kirchhoff_ (1859) on the relation between absorption spectra and the
spectra of luminous incandescent gases. It had already been observed
long before (by Fraunhofer, Foucault, Ångstrom) that the bright spectrum
of the sodium flame gives two bright lines which are in exactly the
same position as two black lines known as D in the solar spectrum,
which evidently belong to an absorption spectrum. When Kirchhoff caused
diffused sunlight to fall upon the slit of a spectroscope, and placed
a sodium flame before it, a perfect superposition was observed--the
bright sodium lines completely covered the black lines D of the solar
spectrum. When further the continuous spectrum of a Drummond light
showed the black line D on placing a sodium flame between it and the
slit of the spectroscope--that is, when the Fraunhofer line of the solar
spectrum was artificially produced--then there was no doubt that its
appearance in the solar spectrum was due to the light passing somewhere
through incandescent vapours of sodium. Hence a new theory of _reversed
spectra_[29] arose--that is, of the relation between the waves of
light emitted and absorbed by a substance under given conditions of
temperature; this is expressed by Kirchhoff's law, discovered by a
careful analysis of the phenomena. This law may be formulated in an
elementary way as follows: At a given temperature the relation between
the intensity of the light emitted (of a definite wave-length) and
the absorptive capacity with respect to the same colour (of the same
wave-length) is a constant quantity.[30] As a black dull surface emits
and also absorbs a considerable quantity of heat rays whilst a polished
metallic surface both absorbs and emits but few, so a flame coloured
by sodium emits a considerable quantity of yellow rays of a definite
refrangibility, and has the property of absorbing a considerable
quantity of the rays of the same refractive index. In general, the
medium which emits definite rays also absorbs them.
[28] The method of observing absorption spectra consists in taking a
continuous spectrum of white light (one which does not show either
dark lines or particularly bright luminous bands--for instance,
the light of a candle, lamp, or other source). The collimator
(that is, the tube with the slit) is directed towards this light,
and then all the colours of the spectrum are visible in the ocular
tube. A transparent absorptive medium--for instance, a solution or
tube containing a gas--is then placed between the source of light
and the apparatus (or anywhere inside the apparatus itself in the
path of the rays). In this case either the entire spectrum is
uniformly fainter, or absorption bands appear on the bright field
of the continuous spectrum in definite positions along it. These
bands have different lengths and positions, and distinctness and
intensity of absorption, according to the properties of the
absorptive medium. Like the luminous spectra given by incandescent
gases and vapours, the absorption spectra of a number of
substances have already been studied, and some with great
precision--as, for example, the spectrum of the brown vapours of
nitrogen dioxide by Hasselberg (at Pulkowa), the spectra of
colouring matters (Eder and others), especially of those applied
to orthochromatic photography, the spectra of blood, chlorophyll
(the green constituent of leaves), and other similar substances,
all the more carefully as by the aid of their spectra the presence
of these substances may be discovered in small quantities (even in
microscopical quantities, by the aid of special appliances on the
microscope), and the changes they undergo investigated.
[Illustration: FIG. 74.--Absorption spectra of nitrogen dioxide
and iodine.]
The absorption spectra, obtained at the ordinary temperature and
proper to substances in all physical states, offer a most
extensive but as yet little studied field, both for the general
theory of spectroscopy, and for gaining an insight into the
structure of substances. The investigation of colouring matters
has already shown that in certain cases a definite change of
composition and structure entails not only a definite change of
the colours but also a displacement of the absorption bands by a
definite number of wave-lengths.
[29] A number of methods have been invented to demonstrate the
reversibility of spectra; among these methods we will cite two
which are very easily carried out. In Bunsen's method sodium
chloride is put into an apparatus for evolving hydrogen (the spray
of the salt is then carried off by the hydrogen and colours the
flame with the yellow sodium colour), and the hydrogen is ignited
in two burners--in one large one with a wide flame giving a bright
yellow sodium light, and in another with a small fine orifice
whose flame is pale: this flame will throw a dark patch on the
large bright flame. In Ladoffsky's method the front tube (p. 561)
is unscrewed from a spectroscope directed towards the light of a
lamp (a continuous spectrum), and the flame of a spirit lamp
coloured by a small quantity of NaCl is placed between the tube
and the prism; a black band corresponding to sodium will then be
seen on looking through the ocular tube. This experiment is always
successful if only there be the requisite relation between the
strength of light of the two lamps.
[30] The absorptive capacity is the relation between the intensity
of the light (of a given wave-length) falling upon and retained by
a substance. Bunsen and Roscoe showed by direct experiment that
this ratio is a constant quantity for every substance. If _A_
stand for this ratio for a given substance at a given
temperature--for instance, for a flame coloured by sodium--and _E_
be the intensity of the light of the same wave-length emitted at
the same temperature by the same substance, then Kirchhoff's law,
the explanation and deduction of which must be looked for in
text-books of physics, states that the fraction _A/E_ is a
constant quantity depending on the nature of a substance (as _A_
depends on it) and determined by the temperature and wave-length.
[Illustration: FIG. 75.--Bright spectra of copper compounds.]
Thus the bright spectral rays characteristic of a given metal may
be reversed--that is, converted into dark lines--by passing light
which gives a continuous spectrum through a space containing the
heated vapours of the given metal. A similar phenomenon to that thus
artificially produced is observed in sunlight, which shows dark lines
characteristic of known metals--that is, the Fraunhofer lines form
an absorption spectrum or depend on a reversed spectrum; it being
presupposed that the sun itself, like all known sources of artificial
light, gives a continuous spectrum without Fraunhofer lines.[31] We
must imagine that the sun, owing to the high temperature which is
proper to it, emits a brilliant light which gives a continuous spectrum,
and that this light, before reaching our eyes, passes through a space
full of the vapours of different metals and their compounds. As the
earth's atmosphere[32] contains very little, or no, metallic vapours,
and as they cannot be supposed to exist in the celestial space,[32 bis]
the only place in which the existence of such vapours can be admitted
is in the _atmosphere surrounding the sun itself_. As the cause of
the sun's luminosity must be looked for in its high temperature, the
existence of an atmosphere containing metallic vapours is readily
understood, because at that high temperature such metals as sodium,
and even iron, are separated from their compounds and converted into
vapour. The sun must be imagined as surrounded by an atmosphere of
incandescent vaporous and gaseous matter,[33] including those elements
whose reversed spectra correspond with the Fraunhofer lines--namely,
sodium, iron, hydrogen, lithium, calcium, magnesium, &c. Thus in
spectrum analysis we find a means of determining the composition of
the inaccessible heavenly luminaries, and much has been done in this
respect since Kirchhoff's theory was formulated. By observations on the
spectra of many heavenly bodies, changes have been discovered going on
in them,[34] and many of the elements known to us have been found with
certainty in them.[35] From this it must be concluded that the same
elements which exist on the earth occur throughout the whole universe,
and that at that degree of heat which is proper to the sun those simple
substances which we accept as the elements in chemistry are still
undecomposed and remain unchanged. A high temperature forms one of those
conditions under which compounds most easily decompose; and if sodium
or a similar element were a compound, in all probability it would be
decomposed into component parts at the high temperature of the sun. This
may indeed be concluded from the fact that in ordinary spectroscopic
experiments the spectra obtained often belong to the metals and not
to the compounds taken; this depends on the decomposition of these
compounds in the heat of the flame. If common salt be introduced
into the flame of a gas-burner, a portion of it is decomposed, first
forming, in all probability, with water, hydrochloric acid and sodium
hydroxide, and the latter then becoming partially decomposed by the
hydrocarbons, giving metallic sodium, whose incandescent vapour emits
light of a definite refrangibility. This conclusion is arrived at from
the following experiment:--If hydrochloric acid gas be introduced into
a flame coloured by sodium it is observed that the sodium spectrum
disappears, owing to the fact that metallic sodium cannot remain in the
flame in the presence of an excess of hydrochloric acid. The same thing
takes place on the addition of sal-ammoniac, which in the heat of the
flame gives hydrochloric acid. If a porcelain tube containing sodium
chloride (or sodium hydroxide or carbonate), and closed at both ends
by glass plates, be so powerfully heated that the salt volatilises,
then the sodium spectrum is not observable; but if the salt be replaced
by sodium, then either the bright line or the absorption spectra is
obtained, according to whether the light emitted by the incandescent
vapour be observed, or light passing through the tube. Thus the above
spectrum is not given by sodium chloride or other sodium compound, but
is proper to the metal sodium itself. This is also the case with other
analogous metals. The chlorides and other halogen _compounds_ of barium,
calcium, copper, &c., give independent spectra which differ from those
of the metals. If barium chloride be introduced into a flame, it gives
a mixed spectrum belonging to metallic barium and barium chloride. If
besides barium chloride, hydrochloric acid or sal-ammoniac be introduced
into the flame, then the spectrum of the metal disappears, and that of
the chloride remains, which differs distinctly from the spectrum of
barium fluoride, barium bromide, or barium iodide. A certain common
resemblance and certain common lines are observed in the spectra of
two different compounds of one and the same element obtained in the
above-described manner, and also in the spectrum of the metal, but they
all have their peculiarities. The independent spectra of the compounds
of copper are easily observed (fig. 75). Thus certain compounds which
exist in a state of vapour, and are luminous at a high temperature,
give their independent spectra. In the majority of cases the spectra of
compounds are composed of indistinct luminous lines and complete bright
bands, whilst metallic elements generally give a few clearly-defined
spectral lines.[36] There is no reason for supposing that the spectrum
of a compound is equal to the sum of the spectra of its elements--that
is, _every compound_ which is not decomposed by heat _has its own
proper spectrum_. This is best proved by absorption spectra, which are
essentially only reversed spectra observed at low temperatures. If every
salt of sodium, lithium, and potassium gives one and the same spectrum,
this must be ascribed to the presence in the flame of the free metals
liberated by the decomposition of their salts. Therefore _the phenomena
of the spectrum are determined by molecules, and not by atoms_--that
is, the molecules of the metal sodium, and not its atoms, produce those
particular vibrations which determine the spectrum of a sodium salt.
Where there is no free metallic sodium there is no sodium spectrum.
[31] Heated metals begin to emit light (only visible in the dark) at
about 420° (varying with the metal). On further heating, solids
first emit red, then yellow, and lastly white light. Compressed or
heavy gases (_see_ Chapter III., Note 44), when strongly heated,
also emit white light. Heated liquids (for example, molten steel
or platinum) also give a white compound light. This is readily
understood. In a dense mass of matter the collisions of the
molecules and atoms are so frequent that waves of only a few
definite lengths cannot appear; the reverse is possible in
rarefied gases or vapours.
[32] Brewster, as is mentioned above, first distinguished the
atmospheric, cosmical Fraunhofer lines from the solar lines.
Janssen showed that the spectrum of the atmosphere contains lines
which depend on the absorption produced by aqueous vapour.
Egoreff, Olszewski, Janssen, and Liveing and Dewar showed by a
series of experiments that the oxygen of the atmosphere gives rise
to certain lines of the solar spectrum, especially the line A.
Liveing and Dewar took a layer of 165 c.m. of oxygen compressed
under a pressure of 85 atmospheres, and determined its absorption
spectrum, and found that, besides the Fraunhofer lines A and B, it
contained the following groups: 630-622, 581-568, 535, 480-475.
The same lines were found for liquid oxygen.
[32 bis] If the material of the whole heavenly space formed the
absorbent medium, the spectra of the stars would be the same as
the solar spectrum; but Huyghens, Lockyer, and others showed not
only that this is the case for only a few stars, but that the
majority of stars give spectra of a different character with dark
and bright lines and bands.
[33] Eruptions, like our volcanic eruptions, but on an incomparably
larger scale, are of frequent occurrence on the sun. They are seen
as protuberances visible during a total eclipse of the sun, in the
form of vaporous masses on the edge of the solar disc and emitting
a faint light. These protuberances of the sun are now observed at
all times by means of the spectroscope (Lockyer's method), because
they contain luminous vapours (giving bright lines) of hydrogen
and other elements.
[34] The great interest and vastness of astro-physical observations
concerning the sun, comets, stars, nebulæ, &c., render this new
province of natural science very important, and necessitate
referring the reader to special works on the subject.
The most important astro-physical data since the time of Kellner
are those referring to the _displacement_ of the lines of the
spectrum. Just as a musical note changes its pitch with the
approach or withdrawal of the resonant object or the ear, so the
pitch of the luminous note or wave-length of the light varies if
the luminous (or absorbent) vapour and the earth from which we
observe it approach or recede from each other; this expresses
itself in a visible displacement of the spectral lines. The solar
eruptions even give broken lines in the spectrum, because the
rapidly moving eruptive masses of vapour and gases either travel
in the direction of the eye or fall back towards the sun. As the
earth travels with the solar system among the stars, so it is
possible to determine the direction and velocity with which the
sun travels in space by the displacement of the spectral lines and
light of the stars. The changes proceeding on the sun in its mass,
which must be pronounced as vaporous, and in its atmosphere, are
now studied by means of the spectroscope. For this purpose, many
special astro-physical observatories now exist where these
investigations are carried on.
We may remark that if the observer or luminous object moves with a
velocity ±_v_, the ray, whose wave-length is [Greek: l], has an
apparent wave-length [Greek: l](_n_±_v_)/_n_, where _n_ is the
velocity of light. Thus Tolon, Huyghens, and others proved that
the star Aldebaran approaches the solar system with a velocity of
30 kilometres per second, while Arcturus is receding with a
velocity of 45 kilometres. The majority of stars give a distinct
hydrogen spectrum, besides which nebulæ also give the spectrum of
nitrogen. Lockyer classes the stars from their spectra, according
to their period of formation, showing that some stars are in a
period of increasing temperature (of formation or aggregation),
whilst others are in a period of cooling. Altogether, in the
astro-physical investigation of the spectra of heavenly bodies we
find one of the most interesting subjects of recent science.
[35] Spectrum analysis has proved the indubitable existence in the sun
and stars of a number of elements known in chemistry. Huyghens,
Secchi, Lockyer, and others have furnished a large amount of
material upon this subject. A compilation of existing information
on it has been given by Prof. S. A. Kleiber, in the Journal of the
Russian Physico-chemical Society for 1885 (vol. xviii. p. 146).
Besides which, a peculiar element called helium has been
discovered, which is characterised by a line (whose wave-length is
587·5, situated near D), which is seen very brightly in the
projections (protuberances) and spots of the sun, but which does
not belong to any known element, and is not reproducible as a
reversed, dark line. This may be a right conclusion--that is to
say, it is possible that an element may be discovered to which the
spectrum of helium corresponds--but it may be that the helium line
belongs to one of the known elements, because spectra vary in the
brilliancy and position of their lines with changes of temperature
and pressure. Thus, for instance, Lockyer could only see the line
423, at the very end of the calcium spectrum, at comparatively low
temperatures, whilst the lines 397 and 393 appear at a higher
temperature, and at a still higher temperature the line 423
becomes quite invisible.
[36] Spectroscopic observations are still further complicated by the
fact that one and the same substance gives different spectra at
different temperatures. This is especially the case with gases
whose spectra are obtained by an electric discharge in tubes.
Plücker, Wüllner, Schuster, and others showed that at low
temperatures and pressures the spectra of iodine, sulphur,
nitrogen, oxygen, &c. are quite different from the spectra of the
same elements at high temperatures and pressures. This may either
depend on the fact that the elements change their molecular
structure with a change of temperature, just as ozone is converted
into oxygen (for instance, from N_{2} molecules are obtained
containing only one atom of nitrogen), or else it may be because
at low temperature certain rays have a greater relative intensity
than those which appear at higher temperatures. If we suppose that
the molecules of a gas are in continual motion, with a velocity
dependent on the temperature, then it must be admitted that they
often strike against each other and rebound, and thus communicate
peculiar motions to each other and the supposed ether, which
express themselves in luminiferous phenomena. A rise of the
temperature or an increase in the density of a gas must have an
influence on the collision of its molecules and luminiferous
motions thus produced, and this may be the cause of the difference
of the spectra under these circumstances. It has been shown by
direct experiment that gases compressed by pressure, when the
collision of the molecules must be frequent and varied, exhibit a
more complex spectrum on the passage of an electric spark than
rarefied gases, and that even a continuous spectrum appears. In
order to show the variability of the spectrum according to the
circumstances under which it proceeds, it may be mentioned that
potassium sulphate fused on a platinum wire gives, on the passage
of a series of sparks, a distinct system of lines, 583-578, whilst
when a series of sparks is passed through a solution of this salt
this system of lines is faint, and when Roscoe and Schuster
observed the absorption spectrum of the vapour of metallic
potassium (which is green) they remarked a number of lines of the
same intensity as the above system in the red, orange, and yellow
portions.
[Illustration: FIG. 76.--Method of showing the spectrum of
substances in solution.]
The spectra of solutions are best observed by means of Lecoq de
Boisbaudran's arrangement, shown in fig. 76. A bent capillary
tube, D F, inside which a platinum wire, A _a_ (from 0·3 to 0·5
mm. in diameter) is fused, is immersed in a narrow cylinder, C (in
which it is firmly held by a cork). The projecting end, _a_, of
the wire is covered by a fine capillary tube, _d_, which extends
1-2 mm. beyond the wire. Another straight capillary tube, E, with
a platinum wire, B _b_, about 1 mm. in diameter (a finer wire soon
becomes hot), is held (by a cork or in a stand) above the end of
the tube, D. If the wire A be now connected with the positive, and
the wire B with the negative terminal of a Ruhmkorff's coil (if
the wires be connected in the opposite order, the spectrum of air
is obtained), a series of sparks rapidly following each other
appear between _a_ and _b_, and their light may be examined by
placing the apparatus in front of the slit of a spectroscope. The
variations to which a spectrum is liable may easily be observed by
increasing the distance between the wires, altering the direction
of the current or strength of the solution, &c.
_Spectrum analysis_ has not only endowed science with a knowledge of
the composition of distant heavenly bodies (of the sun, stars, nebulæ,
comets, &c.), but has also given a new _method_ for studying the matter
of the earth's surface. With its help Bunsen discovered two new elements
belonging to the group of the alkali metals, and thallium, indium, and
gallium were afterwards discovered by the same means. The spectroscope
is employed in the study of rare metals (which in solution often give
distinct absorption spectra), of dyes, and of many organic substances,
&c.[37] With respect to the metals which are analogous to sodium, they
all give similar very volatile salts and such very characteristic
spectra that the least traces of them[38] are discovered with great
ease by means of the spectroscope. For instance, _lithium_ gives a very
brilliant red coloration to a flame and a very bright red spectral line
(wave-length, 670 millionths mm.), which indicates the presence of this
metal in admixture with compounds of other alkali metals.
[37] The importance of the spectroscope for the purpose of chemical
research was already shown by Gladstone in 1856, but it did not
become an accessory to the laboratory until after the discoveries
of Kirchhoff and Bunsen. It may be hoped that in time
spectroscopic researches will meet certain wants of the
theoretical (philosophical) side of chemistry, but as yet all that
has been done in this respect can only be regarded as attempts
which have not yet led to any trustworthy conclusions. Thus many
investigators, by collating the wave-lengths of all the light
vibrations excited by a given element, endeavour to find the law
governing their mutual relations; others (especially Hartley and
Ciamician), by comparing the spectra of analogous elements (for
instance, chlorine, bromine, and iodine), have succeeded in
noticing definite features of resemblance in them, whilst others
(Grünwald) search for relations between the spectra of compounds
and their component elements, &c.; but--owing to the multiplicity
of the spectral lines proper to many elements, and (especially in
the ultra-red and ultra-violet ends of the spectrum) the existence
of lines which are undistinguishable owing to their faintness, and
also owing to the comparative novelty of spectroscopic
research--this subject cannot be considered as in any way
perfected. Nevertheless, in certain instances there is evidently
some relationship between the wave-lengths of all the spectral
lines formed by a given element. Thus, in the hydrogen spectrum
the wave-length = 364·542 _m_^2/(_m_^{2}-4), if _m_ varies as a
series of whole numbers from 3 to 15 (Walmer, Hagebach, and
others). For example, when _m_ = 3, the wave-length of one of the
brightest lines of the hydrogen spectrum is obtained (656·2), when
_m_ = 7, one of the visible violet lines (396·8), and when _m_ is
greater than 9, the ultra-violet lines of the hydrogen spectrum.
[38] In order to show the degree of sensitiveness of spectroscopic
reactions the following observation of Dr. Bence Jones may be
cited: If a solution of 3 grains of a lithium salt be injected
under the skin of a guinea-pig, after the lapse of four minutes,
lithium can be discovered in the bile and liquids of the eye, and,
after ten minutes, in all parts of the animal.
[Illustration: FIG. 77.--Preparation of lithium by the action of a
galvanic current on fused lithium chloride.]
_Lithium_, Li, is, like potassium and sodium, somewhat widely spread
in siliceous rocks, but only occurs in small quantities and as mere
traces in considerable masses of potassium and sodium salts. Only a
very few rather rare minerals contain more than traces of it,[39] for
example, spodumene and lithia mica. Many compounds of lithium are in all
respects closely analogous to the corresponding compounds of sodium
and potassium; but the _carbonate_ is sparingly soluble in cold water,
which fact is taken advantage of for separating lithium from potassium
and sodium. This salt, Li_{2}CO_{3}, is easily converted into the other
compounds of lithium. Thus, for instance, the lithium hydroxide, LiHO,
is obtained in exactly the same way as caustic soda, by the action of
lime on the carbonate, and it is soluble in water and crystallises
(from its solution in alcohol) as LiHO,H_{2}O. Metallic _lithium_ is
obtained by the action of a galvanic current on fused lithium chloride;
for this purpose a cast-iron crucible, furnished with a stout cover,
is filled with lithium chloride, heated until the latter fuses, and a
strong galvanic current is then passed through the molten mass. The
positive pole (fig. 77) consists of a dense carbon rod C (surrounded
by a porcelain tube P fixed in an iron tube BB), and the negative pole
of an iron wire, on which the metal is deposited after the current has
passed through the molten mass for a certain length of time. Chlorine
is evolved at the positive pole. When a somewhat considerable quantity
of the metal has accumulated on the wire it is withdrawn, the metal is
collected from it, and the experiment is then carried on as
before.[39 bis] Lithium is the lightest of all metals, its specific
gravity is 0·59, owing to which it floats even on naphtha; it melts at
180°, but does not volatilise at a red heat. Its appearance recalls that
of sodium, and, like it, it has a yellow tint. At 200° it burns in air
with a very bright flame, forming lithium oxide. In decomposing water it
does not ignite the hydrogen. The characteristic test for lithium
compounds is the _red coloration_ which they impart to a colourless
flame.[40]
[39] Thus _spodumene_ contains up to 6 p.c. of lithium oxide, and
_petolite_, and _lepidolite_ or lithia mica, about 3 p.c. of
lithium oxide. This mica is met with in certain granites in a
somewhat considerable quantity, and is therefore most frequently
employed for the preparation of lithium compounds. The treatment
of lepidolite is carried on on a large scale, because certain
salts of lithium are employed in medicine as a remedy for certain
diseases (stone, gouty affections), as they have the power of
dissolving the insoluble uric acid which is then deposited.
Lepidolite, which is unacted on by acids in its natural state,
decomposes under the action of strong hydrochloric acid after it
has been fused. After being subjected to the action of the
hydrochloric acid for several hours all the silica is obtained in
an insoluble form, whilst the metallic oxides pass into solution
as chlorides. This solution is mixed with nitric acid to convert
the ferrous salts into ferric, and sodium carbonate is then added
until the liquid becomes neutral, by which means a precipitate is
formed of the oxides of iron, alumina, magnesia, &c., as insoluble
oxides and carbonates. The solution (with an excess of water) then
contains the chlorides of the alkaline metals KCl, NaCl, LiCl,
which do not give a precipitate with sodium carbonate in a dilute
solution. It is then evaporated, and a strong solution of sodium
carbonate added. This precipitates lithium carbonate, which,
although soluble in water, is much less so than sodium carbonate,
and therefore the latter precipitates lithium from strong
solutions as carbonate, 2LiCl + Na_{2}CO_{3} = 2NaCl +
Li_{2}CO_{3}. _Lithium carbonate_, which resembles sodium
carbonate in many respects, is a substance which is very slightly
soluble in cold water and is only moderately soluble in boiling
water. In this respect lithium forms a transition between the
metals of the alkalis and other metals, especially those of the
alkaline earths (magnesium, barium), whose carbonates are only
sparingly soluble. Oxide of lithium, Li_{2}O, may be obtained by
heating lithium carbonate with charcoal. Lithium oxide in
dissolving gives (per gram-molecule) 26,000 heat units; but the
combination of Li_{2} with O evolves 140,000 calories--that is,
more than Na_{2}O (100,000 calories) and K_{2}O (97,000 calories),
as shown by Beketoff (1887). Oeuvrard (1892) heated lithium to
redness in nitrogen, and observed the absorption of N and
formation of Li_{3}N, like Na_{3}N (_see_ Chapter XII. Note 50).
LiCl, LiBr, and LiI form crystallo-hydrates with H_{2}O, 2H_{2}O,
and 3H_{2}O. As a rule, LiBr,2H_{2}O crystallises out, but
Bogorodsky (1894) showed that a solution containing LiBr +
3·7H_{2}O, cooled to -62°, separates out crystals LiBr,3H_{2}O,
which decompose at +4° with the separation of H_{2}O. LiF is but
slightly soluble (in 800 parts) in water (and still less so in a
solution of NH_{4}F).
[39 bis] Guntz (1893) recommends adding KCl to the LiCl in preparing Li
by this method, and to act with a current of 10 ampères at 20
volts, and not to heat above 450°, so as to avoid the formation of
Li_{2}Cl.
[40] In determining the presence of lithium in a given compound, it is
best to treat the material under investigation with acid (in the
case of mineral silicon compounds hydrofluoric acid must be
taken), and to treat the residue with sulphuric acid, evaporate to
dryness, and extract with alcohol, which dissolves a certain
amount of the lithium sulphate. It is easy to discover lithium in
such an alcoholic solution by means of the coloration imparted to
the flame on burning it, and in case of doubt by investigating its
light in a spectroscope, because lithium gives a red line, which
is very characteristic and is found as a dark line in the solar
spectrum. Lithium was first discovered in 1817 in petolite by
Arfvedson.
Bunsen in 1860 tried to determine by means of the spectroscope whether
any other as yet unknown metals might not occur in different natural
products together with lithium, potassium, and sodium, and he soon
discovered two new alkali metals showing independent spectra. They are
named after the characteristic coloration which they impart to the
flame. One which gives a red and violet band is named _rubidium_, from
_rubidius_ (dark red), and the other is called _cæsium_, because it
colours a pale flame sky blue, which depends on its containing bright
blue rays, which appear in the spectrum of cæsium as two blue bands
(table on p. 565). Both metals accompany sodium, potassium, and lithium,
but in small quantities; rubidium occurs in larger quantity than cæsium.
The amount of the oxides of cæsium and rubidium in lepidolite does not
generally exceed one-half per cent. Rubidium has also been found in the
ashes of many plants, while the Stassfurt carnallite (the mother-liquor
obtained after having been treated for KCl) forms an abundant source
for rubidium and also partly for cæsium. Rubidium also occurs, although
in very small quantities, in the majority of mineral waters. In a very
few cases cæsium is not accompanied by rubidium; thus, in a certain
granite on the Isle of Elba, cæsium has been discovered, but not
rubidium. This granite contains a very rare mineral called _pollux_,
which contains as much as 34 per cent. of cæsium oxide. Guided by the
spectroscope, and aided by the fact that the double salts of platinic
chloride and rubidium and cæsium chlorides are still less soluble in
water than the corresponding potassium salt, K_{2}PtCl_{6},[41] Bunsen
succeeded in separating both metals from each other and from potassium,
and demonstrated the great resemblance they bear to each other. The
isolated metals,[42] rubidium and cæsium, have respectively the specific
gravities 1·52 and 2·366, and melting points 39° and 27° as N. N.
Beketoff showed (1894), he having obtained cæsium by heating CsAlO_{2}
with Mg([42 bis]).
[41] The salts of the majority of metals are precipitated as carbonates
on the addition of ammonium carbonate--for instance, the salts of
calcium, iron, &c. The alkalis whose carbonates are soluble are
not, however, precipitated in this case. On evaporating the
resultant solution and igniting the residue (to remove the
ammonium salts), we obtain salts of the alkali metals. They may he
separated by adding hydrochloric acid together with a solution of
platinic chloride. The chlorides of lithium and sodium give easily
soluble double salts with platinic chloride, whilst the chlorides
of potassium, rubidium, and cæsium form double salts which are
sparingly soluble. A hundred parts of water at 0° dissolve 0·74
part of the potassium platinochloride; the corresponding rubidium
platinochloride is only dissolved to the amount of 0·134 part, and
the cæsium salt, 0·024 part; at 100° 5·13 parts of potassium
platinochloride, K_{2}PtCl_{6}, are dissolved, 0·634 part of
rubidium platinochloride, and 0·177 part of cæsium
platinochloride. From this it is clear how the salts of rubidium
and cæsium may be isolated. The separation of cæsium from rubidium
by this method is very tedious. It can be better effected by
taking advantage of the difference of the solubility of their
carbonates in alcohol; cæsium carbonate, Cs_{2}CO_{3}, is soluble
in alcohol, whilst the corresponding salts of rubidium and
potassium are almost insoluble. Setterberg separated these metals
as alums, but the best method, that given by Scharples, is founded
on the fact that from a mixture of the chlorides of potassium,
sodium, cæsium, and rubidium in the presence of hydrochloric acid,
stannic chloride precipitates a double salt of cæsium, which is
very slightly soluble. The salts of Rb and Cs are closely
analogous to those of potassium.
[42] Bunsen obtained rubidium by distilling a mixture of the tartrate
with soot, and Beketoff (1888) by heating the hydroxide with
aluminium, 2RbHO + Al = RbAlO_{2} + H_{2} + Rb. By the action of
85 grams of rubidium on water, 94,000 heat units are evolved.
Setterberg obtained cæsium (1882) by the electrolysis of a fused
mixture of cyanide of cæsium and of barium. Winkler (1890) showed
that metallic magnesium reduces the hydrates and carbonates of Rb
and Cs like the other alkaline metals. N. N. Beketoff obtained
them with aluminium (see following note).
[42 bis] Beketoff (1888) showed that metallic aluminium reduces the
hydrates of the alkaline metals at a red heat (they should be
perfectly dry) with the formation of aluminates (Chapter XVII.),
RAlO_{2}--for example, 2KHO + Al = KAlO_{2} + K + H_{2}. It is
evident that in this case only half of the alkaline metal is
obtained free. On the other hand, K. Winkler (1889) showed that
magnesium powder is also able to reduce the alkaline metals from
their hydrates and carbonates. N. N. Beketoff and Tscherbacheff
(1894) prepared cæsium upon this principle by heating its
aluminate CsAlO_{2} with magnesium powder. In this case aluminate
of magnesium is formed, and the whole of the cæsium is obtained as
metal: 2CsAlO_{2} + Mg = MgOAl_{2}O_{5} + 2Cs. A certain excess of
alumina was taken (in order to obtain a less hygroscopic mass of
aluminate), and magnesium powder (in order to decompose the last
traces of water); the CsAlO_{2} was prepared by the precipitation
of cæsium alums by caustic baryta, and evaporating the resultant
solution. We may add that N. N. Beketoff (1887) prepared oxide of
potassium, K_{2}O, by heating the peroxide, KO, in the vapour of
potassium (disengaged from its alloy with silver), and showed that
in dissolving in an excess of water it evolves (for the
above-given molecular weight) 67,400 calories (while 2KHO in
dissolving in water evolves 24,920 cal.; so that K_{2}O + H_{2}O
gives 42,480 cal.), whence (knowing that K_{2} + O+H_{2}O in an
excess of water evolves 164,500) it follows that K_{2} + O evolves
97,100 cal. This quantity is somewhat less than that (100,260
cal.) which corresponds to sodium, and the energy of the action of
potassium upon water is explained by the fact that K_{2}O evolves
more heat than Na_{2}O in combining with water (_see_ Chapter II.
Note 9). Just as hydrogen displaces half the Na from Na_{2}O
forming NaHO, so also N. N. Beketoff found from experiment and
thermochemical reasonings that hydrogen displaces half the
potassium from K_{2}O forming KHO and evolving 7,190 calories.
Oxide of lithium, Li_{2}O, which is easily formed by igniting
Li_{2}CO_{3} with carbon (when Li_{2}O + 2CO is formed),
disengages 26,000 cals. with an excess of water, while the
reaction Li_{2} + O gives 114,000 cals. and the reaction Li_{2} +
H_{2}O gives only 13,000 cals., and metallic lithium cannot be
liberated from oxide of lithium with hydrogen (nor with carbon).
Thus in the series Li, Na, K, the formation of R_{2}O gives most
heat with Li and least with K, while the formation of RCl evolves
most heat with K (105,000 cals.) and least of all with Li (93,500
cals.). Rubidium, in forming Rb_{2}O, gives 94,000 cals.
(Beketoff). Cæsium, in acting upon an excess of water, evolves
51,500 cals., and the reaction Cs_{2} + O evolves about 100,000
cals.--_i.e._ more than K and Rb, and almost as much as Na--and
oxide of cæsium reacts with hydrogen (according to the equation
Cs_{2}O + H = CsHO + Cs) more easily than any of the oxides of the
alkali metals, and this reaction takes place at the ordinary
temperature (the hydrogen is absorbed), as Beketoff showed (1893).
He also obtained a mixed oxide, AgCsO, which was easily formed in
the presence of silver, and absorbed hydrogen with the formation
of CsHO.
Judging by the properties of the free metals, and of their corresponding
and even very complex compounds, lithium, sodium, potassium, rubidium,
and cæsium present an indubitable chemical resemblance. The fact that
the metals easily decompose water, and that their hydroxides RHO and
carbonates R_{2}CO_{3} are soluble in water, whilst the hydroxides
and carbonates of nearly all other metals are insoluble, shows that
these metals form a natural group of _alkali metals_. The halogens and
the alkali metals form, by their character, the two extremes of the
elements. Many of the other elements are metals approaching the alkali
metals, both in their capacity of forming salts and in not forming acid
compounds, but are not so energetic as the alkali metals, that is, they
form less energetic bases. Such are the common metals, silver, iron,
copper, &c. Some other elements, in the character of their compounds,
approach the halogens, and, like them, combine with hydrogen, but these
compounds do not show the energetic property of the halogen acids; in
a free state they easily combine with metals, but they do not then
form such saline compounds as the halogens do--in a word, the halogen
properties are less sharply defined in them than in the halogens
themselves. Sulphur, phosphorus, arsenic, &c. belong to this order of
elements. The clearest distinction of the properties of the halogens
and alkali metals is expressed in the fact that the former give acids
and do not form bases, whilst the latter, on the contrary, only give
bases. The first are true _acid elements_, the latter clearly-defined
_basic or metallic elements_. On combining together, the halogens form,
in a chemical sense, unstable compounds, and the alkali metals alloys
in which the character of the metals remains unaltered, just as in the
compound ICl the character of the halogens remains undisguised; thus
both classes of elements on combining with members of their own class
form non-characteristic compounds, which have the properties of their
components. On the other hand, the halogens on combining with the alkali
metals form compounds which are, in all respects, stable, and in which
the original characters of the halogens and alkali metals have entirely
disappeared. The formation of such compounds is accompanied by evolution
of a large amount of heat, and by an entire change of both the physical
and chemical properties of the substances originally taken. The alloy of
sodium and potassium, although liquid at the ordinary temperature, is
perfectly metallic, like both its components. The compound of sodium and
chlorine has neither the appearance nor the properties of the original
elements; sodium chloride melts at a higher temperature, and is more
difficultly volatile, than either sodium or chlorine.
With all these qualitative differences there is, however, an important
quantitative _resemblance between the halogens and the alkali metals_.
This resemblance is clearly expressed by stating that both orders of
elements belong to those which are univalent with respect to hydrogen.
It is thus correct to say that both the above-named orders of elements
replace hydrogen atom for atom. Chlorine is able to take the place of
hydrogen by metalepsis, and the alkali metals take the place of hydrogen
in water and acids. As it is possible to consecutively replace every
equivalent of hydrogen in a hydrocarbon by chlorine, so it is possible
in an acid containing several equivalents of hydrogen to replace the
hydrogen consecutively equivalent after equivalent by an alkali metal;
hence an atom of these elements is analogous to an atom of hydrogen,
which is taken, in all cases, as the unit for the comparison of the
other elements. In ammonia, and in water, chlorine and sodium are
able to bring about a direct replacement. According to the law of
substitution, the formation of sodium chloride, NaCl, at once shows the
equivalence of the atoms of the alkali metals and the halogens. The
halogens and hydrogen and the alkali metals combine with such elements
as oxygen, and it is easily proved that in such compounds one atom
of oxygen is able to retain two atoms of the halogens, of hydrogen,
and of the alkali metals. For this purpose it is enough to compare
the compounds KHO, K_{2}O, HClO, and Cl_{2}O, with water. It must not
be forgotten, however, that the halogens give, with oxygen, besides
compounds of the type R_{2}O, higher acid grades of oxidation, which the
alkali metals and hydrogen are not capable of forming. We shall soon
see that these relations are also subject to a special law, showing a
gradual transition of the properties of the elements from the alkali
metals to the halogens.[43]
[43] We may here observe that the halogens, and especially iodine,
may play the part of metals (hence iodine is more easily replaced
by metals than the other halogens, and it approaches nearer to the
metals in its physical properties than the other halogens).
Schützenberger obtained a compound C_{2}H_{3}O(OCl), which he
called chlorine acetate, by acting on acetic anhydride,
(C_{2}H_{3}O)_{2}O, with chlorine monoxide, Cl_{2}O. With iodine
this compound gives off chlorine and forms iodine acetate,
C_{2}H_{3}O(OI), which also is formed by the action of iodine
chloride on sodium acetate, C_{2}H_{3}O(ONa). These compounds are
evidently nothing else than mixed anhydrides of hypochlorous and
hypoiodous acids, or the products of the substitution of hydrogen
in RHO by a halogen (_see_ Chapter XI., Notes 29 and 78 bis). Such
compounds are very unstable, decompose with an explosion when
heated, and are changed by the action of water and of many other
reagents, which is in accordance with the fact that they contain
very closely allied elements, as does Cl_{2}O itself, or ICl or
KNa. By the action of chlorine monoxide on a mixture of iodine and
acetic anhydride, Schützenberger also obtained the compound
I(C_{2}H_{3}O_{2})_{3}, which is analogous to ICl_{3}, because the
group C_{2}H_{3}O_{2} is, like Cl, a halogen, forming salts with
the metals. Similar properties are found in iodosobenzene (Chapter
XI., Note 79).
The atomic weights of the alkali metals, lithium 7, sodium 23, potassium
39, rubidium 85, and cæsium 133, show that here, as in the class of
halogens, the elements may be arranged according to their atomic weights
in order to compare the properties of the analogous compounds of the
members of this group. Thus, for example, the platinochlorides of lithium
and sodium are soluble in water; those of potassium, rubidium, and
cæsium sparingly soluble, and the greater the atomic weight of the metal
the less soluble is the salt. In other cases the reverse is observed--the
greater the atomic weight the more soluble are the corresponding salts.
The variation of properties with the variation in atomic weights even
shows itself in the metals themselves; thus lithium volatilises with
difficulty, whilst sodium is obtained by distillation, potassium
volatilises more easily than sodium, and rubidium and cæsium as we have
seen, are still more volatile.
CHAPTER XIV
THE VALENCY AND SPECIFIC HEAT OF THE METALS. MAGNESIUM. CALCIUM,
STRONTIUM, BARIUM, AND BERYLLIUM
It is easy by investigating the composition of corresponding compounds,
to establish the _equivalent weights_ of the metals compared with
hydrogen--that is, the quantity which replaces one part by weight of
hydrogen. If a metal decomposes acids directly, with the evolution of
hydrogen, the equivalent weight of the metal may be determined by taking
a definite weight of it and measuring the volume of hydrogen evolved by
its action on an excess of acid; it is then easy to calculate the weight
of the hydrogen from its volume.[1] The same result may be arrived at
by determining the composition of the normal salts of the metal; for
instance, by finding the weight of metal which combines with 35·5 parts
of chlorine or 80 parts of bromine.[2] The equivalent of a metal may be
also ascertained by simultaneously (_i.e._ in one circuit) decomposing
an acid and a fused salt of a given metal by an electric current
and determining the relation between the amounts of hydrogen and
metal separated, because, according to Faraday's law, electrolytes
(conductors of the second order) are always decomposed in equivalent
quantities.[2 bis] The equivalent of a metal may even be found by simply
determining the relation between its weight and that of its salt-giving
oxide, as by this we know the quantity of the metal which combines with 8
parts by weight of oxygen, and this will be the equivalent, because 8
parts of oxygen combine with 1 part by weight of hydrogen. One method is
verified by another, and all the processes for the accurate determination
of equivalents require the greatest care to avoid the absorption of
moisture, further oxidation, volatility, and other accidental influences
which affect exact weighings. The description of the methods necessary
for the attainment of exact results belongs to the province of analytical
chemistry.
[1] Under favourable circumstances (by taking all the requisite
precautions), the weight of the equivalent may be accurately
determined by this method. Thus Reynolds and Ramsay (1887)
determined the equivalent of zinc to be 32·7 by this method (from
the average of 29 experiments), whilst by other methods it has been
fixed (by different observers) between 32·55 and 33·95.
The differences in their equivalents may be demonstrated by taking
equal weights of different metals, and collecting the hydrogen
evolved by them (under the action of an acid or alkali).
[2] The most accurate determinations of this kind were carried on by
Stas, and will be described in Chapter XXIV.
[2 bis] The amount of electricity in one coulomb according to the
present nomenclature of electrical units (_see_ Works on Physics
and Electro-technology) disengages 0·00001036 gram of hydrogen,
0·00112 gram of silver, 0·0003263 gram of copper from the salts of
the oxide, and 0·0006526 gram from the salts of the suboxide, &c.
These amounts stand in the same ratio as the equivalents, _i.e._ as
the quantities replaced by one part by weight of hydrogen. The
intimate bond which is becoming more and more marked existing
between the electrolytic and purely chemical relations of
substances (especially in solutions) and the application of
electrolysis to the preparation of numerous substances on a large
scale, together with the employment of electricity for obtaining
high temperatures, &c., makes me regret that the plan and
dimensions of this book, and the impossibility of giving a concise
and objective exposition of the necessary electrical facts, prevent
my entering upon this province of knowledge, although I consider it
my duty to recommend its study to all those who desire to take part
in the further development of our science.
There is only one side of the subject respecting the direct
correlation between thermochemical data and electro-motive force,
which I think right to mention here, as it justifies the general
conception, enunciated by Faraday, that the galvanic current is an
aspect of the transference of chemical motion or reaction along the
conductors.
From experiments conducted by Favre, Thomsen, Garni, Berthelot,
Cheltzoff, and others, upon the amount of heat evolved in a closed
circuit, it follows that the electro-motive force of the current or
its capacity to do a certain work, E, is proportional to the whole
amount of heat, Q, disengaged by the reaction forming the source of
the current. If E be expressed in volts, and Q in thousands of
units of heat referred to equivalent weights, then E = 0·0436Q. For
example in a Daniells battery E = 1·09 both by experiment and
theory, because in it there takes place the decomposition of
CuSO_{4} into Cu + O together with the formation of Zn + O and ZnO
+ SO_{3}Aq, and these reactions correspond to Q = 25·06 thousand
units of heat. So also in all other primary batteries (_e.g._
Bunsen's, Poggendorff's, &c.) and secondary ones (for instance,
those acting according to the reaction Pb + H_{2}SO_{4} + PbO_{2},
as Cheltzoff showed) E = 0·0436Q.
For univalent metals, like those of the alkalis, the weight of the
equivalent is equal to the weight of the atom. For bivalent metals the
atomic weight is equal to the weight of two equivalents, for _n_-valent
metals it is equal to the weight of _n_ equivalents. Thus aluminium,
Al = 27, is trivalent, that is, its equivalent = 9; magnesium, Mg =
24, is bivalent, and its equivalent = 12. Therefore, if potassium or
sodium, or in general a univalent metal, M, give compounds M_{2}O, MHO,
MCl, MNO_{3}, M_{2}SO_{4}, &c., and in general MX, then for bivalent
metals like magnesium or calcium the corresponding compounds will be
MgO, Mg(HO)_{2}, MgCl_{2}, Mg(NO_{3})_{2}, MgSO_{4}, &c., or in general
MX_{2}.
By what are we to be guided in ascribing to some metals univalency and
to others bi-, ter-, quadri-, ... _n_-valency? What obliges us to make
this difference? Why are not all metals given the same valency--for
instance, why is not magnesium considered as univalent? If this be
done, taking Mg = 12 (and not 24 as now), not only is a simplicity
of expression of the composition of all the compounds of magnesium
attained, but we also gain the advantage that their composition will
be the same as those of the corresponding compounds of sodium and
potassium. These combinations were so expressed formerly--why has this
since been changed?
These questions could only be answered after the establishment of the
idea of multiples of the atomic weights as the minimum quantities of
certain elements combining with others to form compounds--in a word,
since the time of the establishment of Avogadro-Gerhardt's law (Chapter
VII.). By taking such an element as arsenic, which has many volatile
compounds, it is easy to determine the density of these compounds, and
therefore to establish their molecular weights, and hence to find the
indubitable atomic weight, exactly as for oxygen, nitrogen, chlorine,
carbon, &c. It appears that As = 75, and its compounds correspond, like
the compounds of nitrogen, with the forms AsX_{3}, and AsX_{5}; for
example, AsH_{3}, AsCl_{3}, AsF_{5}, As_{2}O_{5}, &c. It is evident that
we are here dealing with a metal (or rather element) of two valencies,
which moreover is never univalent, but tri- or quinqui-valent. This
example alone is sufficient for the recognition of the existence of
polyvalent atoms among the metals. And as antimony and bismuth are
closely analogous to arsenic in all their compounds, (just as potassium
is analogous to rubidium and cæsium); so, although very few volatile
compounds of bismuth are known, it was necessary to ascribe to them
formulæ corresponding with those ascribed to arsenic.
As we shall see in describing them, there are also many analogous metals
among the bivalent elements, some of which also give volatile compounds.
For example, zinc, which is itself volatile, gives several volatile
compounds (for instance, zinc ethyl, ZnC_{4}H_{10}, which boils at 118°,
vapour density = 61·3), and in the molecules of all these compounds
there is never less than 65 parts of zinc, which is equivalent to H_{2},
because 65 parts of zinc displace 2 parts by weight of hydrogen; so
that zinc is just such an example of the bivalent metals as oxygen,
whose equivalent = 8 (because H_{2} is replaced by O = 16), is a
representative of the bivalent elements, or as arsenic is of the tri-
and quinqui-valent elements. And, as we shall afterwards see, magnesium
is in many respects closely analogous to zinc, which fact obliges us to
regard magnesium as a bivalent metal.
Such metals as mercury and copper, which are able to give not one but
two bases, are of particular importance for distinguishing univalent and
bivalent metals. Thus copper gives the suboxide Cu_{2}O and the oxide
CuO--that is, the compounds CuX corresponding with the suboxide are
analogous (in the quantitative relations, by their composition) to NaX
or AgX, and the compounds of the oxide CuX_{2}, to MgX_{2}, ZnX_{2}, and
in general to the bivalent metals. It is clear that in such examples we
must make a distinction between atomic weights and equivalents.
In this manner the valency, that is, the number of equivalents entering
into the atom of the metals may in many cases be established by means
of comparatively few volatile metallic compounds, with the aid of a
search into their analogies (concerning which see Chapter XV.). _The law
of specific heats_ discovered by Dulong and Petit has frequently been
applied to the same purpose[3] in the history of chemistry, especially
since the development given to this law by the researches of Regnault,
and since Cannizzaro (1860) showed the agreement between the deductions
of this law and the consequences arising from Avogadro-Gerhardt's law.
[3] The chief means by which we determine the valency of the elements,
or what multiple of the equivalent should be ascribed to the atom,
are: (1) The law of Avogadro-Gerhardt. This method is the most
general and trustworthy, and has already been applied to a great
number of elements. (2) The different grades of oxidation and their
isomorphism or analogy in general; for example, Fe = 56 because the
suboxide (ferrous oxide) is isomorphous with magnesium oxide, &c.,
and the oxide (ferric oxide) contains half as much oxygen again as
the suboxide. Berzelius, Marignac, and others took advantage of
this method for determining the composition of the compounds of
many elements. (3) The specific heat, according to Dulong and
Petit's law. Regnault, and more especially Cannizzaro, used this
method to distinguish univalent from bivalent metals. (4) The
periodic law (_see_ Chapter XV.) has served as a means for the
determination of the atomic weights of cerium, uranium, yttrium,
&c., and more especially of gallium, scandium, and germanium. The
correction of the results of one method by those of others is
generally had recourse to, and is quite necessary, because,
phenomena of dissociation, polymerisation, &c., may complicate the
individual determinations by each method.
It will be well to observe that a number of other methods,
especially from the province of those physical properties which are
clearly dependent on the magnitude of the atom (or equivalent) or
of the molecule, may lead to the same result. I may point out, for
instance, that even the specific gravity of solutions of the
metallic chlorides may serve for this purpose. Thus, if beryllium
he taken as trivalent--that is, if the composition of its chloride
be taken as BeCl_{3} (or a polymeride of it), then the specific
gravity of solutions of beryllium chloride will not fit into the
series of the other metallic chlorides. But by ascribing to it an
atomic weight Be = 7, or taking Be as bivalent, and the composition
of its chloride as BeCl_{2}, we arrive at the general rule given in
Chapter VII., Note 28. Thus W. G. Burdakoff determined in my
laboratory that the specific gravity at 15°/4° of the solution
BeCl_{2} + 200H_{2}O = 1·0138--that is, greater than the
corresponding solution KCl + 200H_{2}O (= 1·0121), and less than
the solution MgCl_{2} + 200H_{2}O (= 1·0203), as would follow from
the magnitude of the molecular weight BeCl_{2} = 80, since KCl =
74·5 and MgCl_{2} = 95.
Dulong and Petit, having determined the specific heat of a number of
solid elementary substances, observed that as the atomic weights of the
elements increase, their specific heats decrease, and that _the product
of the specific heat Q into the atomic weight A is an almost constant
quantity_. This means that to bring different elements into a known
thermal state an equal amount of work is required if atomic quantities
of the elements are taken; that is, the amounts of heat expended in
heating equal quantities by weight of the elements are far from equal,
but are in inverse proportion to the atomic weights. For thermal
changes the atom is a unit; all atoms, notwithstanding the difference
of weight and nature, are equal. This is the simplest expression of the
fact discovered by Dulong and Petit. The specific heat measures that
quantity of heat which is required to raise the temperature of _one
unit of weight_ of a substance by one degree. If the magnitude of the
specific heat of elements be multiplied by the atomic weight, then we
obtain the atomic heat--that is, the amount of heat required to raise
the temperature of the atomic weight of an element by one degree. It
is these products which for the majority of the elements prove to be
approximately, if not quite, identical. A complete identity cannot be
expected, because the specific heat of one and the same substance varies
with the temperature, with its passage from one state into another, and
frequently with even a simple mechanical change of density (for instance
by hammering), not to speak of allotropic changes, &c. We will cite
several figures[4] proving the truth of the conclusions arrived at by
Dulong and Petit with respect to solid elementary bodies.
Li Na Mg P
A = 7 23 24 31
Q = 0·9408 0·2934 0·245 0·202
AQ = 6·59 6·75 5·88 6·26
Fe Cu Zn Br
A = 56 63 65 80
Q = 0·112 0·093 0·093 0·0843
AQ = 6·27 5·86 6·04 6·74
Pd Ag Sn I
A = 106 108 118 127
Q = 0·0592 0·056 0·055 0·541
AQ = 6·28 6·05 6·49 6·87
Pt Au Hg Pb
A = 196 198 200 206
Q = 0·0325 0·0324 0·0333 0·0315
AQ = 6·37 6·41 6·66 6·49