SODIUM
The neutral salt, sodium sulphate, Na_{2}SO_{4}, obtained when a mixture
of sulphuric acid and common salt is strongly heated (Chapter X.),[1]
forms a colourless saline mass consisting of fine crystals, soluble in
water. It is the product of many other double decompositions, sometimes
carried out on a large scale; for example, when ammonium sulphate is
heated with common salt, in which case the sal-ammoniac is volatilised,
&c. A similar decomposition also takes place when, for instance, a
mixture of lead sulphate and common salt is heated; this mixture easily
fuses, and if the temperature be further raised heavy vapours of lead
chloride appear. When the disengagement of these vapours ceases, the
remaining mass, on being treated with water, yields a solution of
sodium sulphate mixed with a solution of undecomposed common salt. A
considerable quantity, however, of the lead sulphate remains unchanged
during this reaction, PbSO_{4} + 2NaCl = PbCl_{2} + Na_{2}SO_{4}, the
vapours will contain lead chloride, and the residue will contain the
mixture of the three remaining salts. The cause and nature of the
reaction are just the same as were pointed out when considering the
action of sulphuric acid upon NaCl. Here too it may be shown that the
double decomposition is determined by the removal of PbCl_{2} from the
sphere of the action of the remaining substances. This is seen from the
fact that sodium sulphate, on being dissolved in water and mixed with a
solution of any lead salt (and even with a solution of lead chloride,
although this latter is but sparingly soluble in water), immediately
gives a white precipitate of lead sulphate. In this case the lead takes
up the elements of sulphuric acid from the sodium sulphate in the
solutions. On heating, the reverse phenomenon is observed. The reaction
in the solution depends upon the insolubility of the lead sulphate, and
the decomposition which takes place on heating is due to the volatility
of the lead chloride. Silver sulphate, Ag_{2}SO_{4}, in solution with
common salt, gives silver chloride, because the latter is insoluble in
water, Ag_{2}SO_{4} + 2NaCl = Na_{2}SO_{4} + 2AgCl. Sodium carbonate,
mixed in solution with the sulphates of iron, copper, manganese,
magnesium, &c., gives in solution sodium sulphate, and in the precipitate
a carbonate of the corresponding metal, because these salts of carbonic
acid are insoluble in water; for instance, MgSO_{4} + Na_{2}CO_{3} =
Na_{2}SO_{4} + MgCO_{3}. In precisely the same way sodium hydroxide acts
on solutions of the majority of the salts of sulphuric acid containing
metals, the hydroxides of which are insoluble in water--for instance,
CuSO_{4} + 2NaHO = Cu(HO)_{2} + Na_{2}SO_{4}. Sulphate of magnesium,
MgSO_{4}, on being mixed in solution with common salt, forms, although
not completely, chloride of magnesium, and sodium sulphate. On cooling
the mixture of such (concentrated) solutions sodium sulphate is
deposited, as was shown in Chapter X. This is made use of for preparing
it on the large scale in works where sea-water is treated. In this case,
on cooling, the reaction 2NaCl + MgSO_{4} = MgCl_{2} + Na_{2}SO_{4} takes
place.
[1] Whilst describing in some detail the properties of sodium
chloride, hydrochloric acid, and sodium sulphate, I wish to impart,
by separate examples, an idea of the properties of saline
substances, but the dimensions of this work and its purpose and aim
do not permit of entering into particulars concerning every salt,
acid, or other substance. The fundamental object of this work--an
account of the characteristics of the elements and an acquaintance
with the forces acting between atoms--has nothing to gain from the
multiplication of the number of as yet ungeneralised properties and
relations.
Thus where sulphates and salts of sodium are in contact, it may be
expected that sodium sulphate will be formed and separated if the
conditions are favourable; for this reason it is not surprising that
sodium sulphate is often found in the native state. Some of the springs
and salt lakes in the steppes beyond the Volga, and in the Caucasus,
contain a considerable quantity of sodium sulphate, and yield it by
simple evaporation of the solutions. Beds of this salt are also met with;
thus at a depth of only 5 feet, about 38 versts to the east of Tiflis, at
the foot of the range of the 'Wolf's mane' (Voltchia griva) mountains, a
deep stratum of very pure Glauber's salt, Na_{2}SO_{4},10H_{2}O, has been
found.[2] A layer two metres thick of the same salt lies at the bottom of
several lakes (an area of about 10 square kilometres) in the Kouban
district near Batalpaschinsk, and here its working has been commenced
(1887). In Spain, near Arangoulz and in many parts of the Western States
of North America, mineral sodium sulphate has likewise been found, and is
already being worked.
[2] Anhydrous (ignited) sodium sulphate, Na_{2}SO_{4}, is known in
trade as 'sulphate' or salt-cake, in mineralogy _thenardite_.
Crystalline decahydrated salt is termed in mineralogy _mirabilite_,
and in trade Glauber's salt. On fusing it, the monohydrate
Na_{2}SO_{4}H_{2}O is obtained, together with a supersaturated
solution.
The methods of obtaining salts by means of double decomposition from
others already prepared are so general, that in describing a given salt
there is no necessity to enumerate the cases hitherto observed of its
being formed through various double decompositions.[3] The possibility
of this occurrence ought to be foreseen according to Berthollet's
doctrine from the properties of the salt in question. On this account it
is important to know the properties of salts; all the more so because
up to the present time those very properties (solubility, formation
of crystallo-hydrates, volatility, &c.) which may be made use of for
separating them from other salts have not been generalised.[4] These
properties as yet remain subjects for investigation, and are rarely
to be foreseen. The crystallo-hydrate of the normal sodium sulphate,
Na_{2}SO_{4},10H_{2}O, very easily parts with water, and may be obtained
in an anhydrous state if it be carefully heated until the weight
remains constant; but if heated further, it partly loses the elements
of sulphuric anhydride. The normal salt fuses at 843° (red heat), and
volatilises to a slight extent when very strongly heated, in which
case it naturally decomposes with the evolution of SO_{3}. At 0° 100
parts of water dissolve 5 parts of the anhydrous salt, at 10° 9 parts,
at 20° 19·4, at 30° 40, and at 34° 55 parts, the same being the case
in the presence of an excess of crystals of Na_{2}SO_{4},10H_{2}O.[5]
At 34° the latter fuses, and the solubility decreases at higher
temperatures.[6] A concentrated solution at 34° has a composition
nearly approaching to Na_{2}SO_{4} + 14H_{2}O, and the decahydrated salt
contains 78·9 of the anhydrous salt combined with 100 parts of water.
From the above figures it is seen that the decahydrated salt cannot
fuse without decomposing,[7] like hydrate of chlorine, Cl_{2},8H_{2}O
(Chapter XI., Note 10). Not only the fused decahydrated salt, but also
the concentrated solution at 34° (not all at once, but gradually),
yields the monohydrated salt, Na_{2}SO_{4},H_{2}O. The heptahydrated
salt, Na_{2}SO_{4},7H_{2}O, also splits up, even at low temperatures,
with the formation of this monohydrated salt, and therefore from 35°
the solubility can be given only for the latter. For 100 parts of water
this is as follows: at 40° 48·8, at 50° 46·7, at 80° 43·7, at 100°
42·5 parts of the anhydrous salt. If the decahydrated salt be fused,
and the solution allowed to cool in the presence of the monohydrated
salt, then at 30° 50·4 parts of anhydrous salt are retained in the
solution, and at 20° 52·8 parts. Hence, with respect to the anhydrous
and monohydrated salts, the solubility is identical, and falls with
increasing temperature, whilst with respect to decahydrated salt, the
solubility rises with increasing temperature. So that if in contact
with a solution of sodium sulphate there are only crystals of that
heptahydrated salt (Chapter I., Note 54), Na_{2}SO_{4},7H_{2}O, which
is formed from saturated solutions, then saturation sets in when the
solution has the following composition per 100 parts of salt: at 0°
19·6, at 10° 30·5, at 20° 44·7, and at 25° 52·9 parts of anhydrous
salt. Above 27° the heptahydrated salt, like the decahydrated
salt at 34°, splits up into the monohydrated salt and a saturated
solution. Thus sodium sulphate has three curves of solubility: one for
Na_{2}SO_{4},7H_{2}O (from 0° to 26°), one for Na_{2}SO_{4},10H_{2}O
(from 0° to 34°), and one for Na_{2}SO_{4},H_{2}O (a descending curve
beginning at 26°), because there are three of these crystallo-hydrates,
and the solubility of a substance only depends upon the particular
condition of that portion of it which has separated from the solution or
is present in excess.[8]
[3] The salts may be obtained not only by methods of substitution of
various kinds, but also by many other combinations. Thus sodium
sulphate may be formed from sodium oxide and sulphuric anhydride,
by oxidising sodium sulphide, Na_{2}S, or sodium sulphite,
Na_{2}SO_{3}, &c. When sodium chloride is heated in a mixture of
the vapours of water, air, and sulphurous anhydride, sodium
sulphate is formed. According to this method (patented by
Hargreaves and Robinson), sodium sulphate, Na_{2}SO_{4}, is
obtained from NaCl without the preliminary manufacture of
H_{2}SO_{4}. Lumps of NaCl pressed into bricks are loosely packed
into a cylinder and subjected, at a red heat, to the action of
steam, air and SO_{2}. Under these conditions, HCl, sulphate, and a
certain amount of unaltered NaCl are obtained. This mixture is
converted into soda by Gossage's process (_see_ Note 15) and may
have some practical value.
[4] Many observations have been made, but little general information
has been obtained from particular cases. In addition to which, the
properties of a given salt are changed by the presence of other
salts. This takes place not only in virtue of mutual decomposition
or formation of double salts capable of separate existence, but is
determined by the influence which some salts exert on others, or by
forces similar to those which act during solution. Here nothing has
been generalised to that extent which would render it possible to
predict without previous investigation, if there be no close
analogy to help us. Let us state one of these numerous cases: 100
parts of water at 20° dissolve 34 parts of potassium nitrate but on
the addition of sodium nitrate the solubility of potassium nitrate
increases to 48 parts in 10 of water (Carnelley and Thomson). In
general, in all cases of which there are accurate observations it
appears that the presence of foreign salts changes the properties
of any given salt.
[5] The information concerning solubility (Chapter I.) is given
according to the determinations of Gay-Lussac, Lovell, and Mulder.
[6] In Chapter I., Note 24, we have already seen that with many other
sulphates the solubility also decreases after a certain temperature
is passed. Gypsum, CaSO_{4},2H_{2}O, lime, and many other compounds
present such a phenomenon. An observation of Tilden's (1884) is
most instructive; he showed that on raising the temperature (in
closed vessels) above 140° the solubility of sodium sulphate again
begins to increase. At 100° 100 parts of water dissolve about 43
parts of anhydrous salt, at 140° 42 parts, at 160° 43 parts, at
180° 44 parts, at 230° 46 parts. According to Étard (1892) the
solubility of 30 parts of Na{2}SO_{4} in 100 of solution (or 43 per
100 of water) corresponds to 80°, and above 240° the solubility
again falls, and very rapidly, so that at 320° the solution
contains 12 per 100 of solution (about 14 per 100 of water) and a
further rise of temperature is followed by a further deposition of
the salt. It is evident that the phenomenon of saturation,
determined by the presence of an excess of the dissolved substance,
is very complex, and therefore that for the theory of solutions
considered as liquid indefinite chemical compounds, many useful
statements can hardly be given.
[7] Already referred to in Chapter I., Note 56.
The example of sodium sulphate is historically very important for
the theory of solutions. Notwithstanding the number of
investigations which have been made, it is still insufficiently
studied, especially from the point of the vapour tension of
solutions and crystallo-hydrates, so that those processes cannot be
applied to it which Guldberg, Roozeboom, Van't Hoff, and others
applied to solutions and crystallo-hydrates. It would also be most
important to investigate the influence of pressure on the various
phenomena corresponding with the combinations of water and sodium
sulphate, because when crystals are separated--for instance, of the
decahydrated salt--an increase of volume takes place, as can be
seen from the following data:--the sp. gr. of the anhydrous salt is
2·66, that of the decahydrated salt = 1·46, but the sp. gr. of
solutions at 15°/4° = 9,992 + 90·2_p_ + 0·35_p_^2 where p
represents the percentage of anhydrous salt in the solution, and
the sp. gr. of water at 4° = 10,000. Hence for solutions containing
20 p.c. of anhydrous salt the sp. gr. = 1·1936; therefore the
volume of 100 grams of this solution = 83·8 c.c., and the volume of
anhydrous salt contained in it is equal to 20/2·66, or = 7·5 c.c.,
and the volume of water = 80·1 c.c. Therefore, the solution, on
decomposing into anhydrous salt and water, increases in volume
(from 83·8 to 87·6); but in the same way 83·8 c.c. of 20 p.c.
solution are formed from (45·4/1·46 =) 31·1 c.c. of the
decahydrated salt, and 54·6 c.c. of water--that is to say, that
during the formation of a solution from 85·7 c.c., 83·8 c.c. are
formed.
[8] From this example it is evident the solution remains unaltered
until from the contact of a solid it becomes either saturated or
supersaturated, crystallisation being determined by the attraction
to a solid, as the phenomenon of supersaturation clearly
demonstrates. This partially explains certain apparently
contradictory determinations of solubility. The best investigated
example of such complex relations is cited in Chapter XIV., Note 50
(for CaCl_{2}).
Thus solutions of sodium sulphate may give crystallo-hydrates of three
kinds on cooling the saturated solution: the unstable heptahydrated salt
is obtained at temperatures below 26°, the decahydrated salt forms under
ordinary conditions at temperatures below 34°, and the monohydrated salt
at temperatures above 34°. Both the latter crystallo-hydrates present a
stable state of equilibrium, and the heptahydrated salt decomposes into
them, probably according to the equation 3Na_{2}SO_{4},7H_{2}0 =
2Na_{2}SO_{4},10H_{2}O + Na_{2}SO_{4},H_{2}O. The ordinary decahydrated
salt is called _Glauber's salt_. All forms of these crystallo-hydrates
lose their water entirely, and give the anhydrous salt when dried over
sulphuric acid.[9]
[9] According to Pickering's experiments (1886), the molecular weight
in grams (that is, 142 grams) of anhydrous sodium sulphate, on
being dissolved in a large mass of water, at 0° absorbs (hence
the-sign)-1,100 heat units, at 10°-700, at 15°-275, at 20° gives
out +25, at 25° +300 calories. For the decahydrated salt,
Na_{2}SO_{4},10H_{2}O, 5°-4,225, 10°-4,000, 15°-3,570, 20°-3,160,
25°-2,775. Hence (just as in Chapter I., Note 56) the heat of the
combination Na_{2}SO_{4},10H_{2}O at 5° = +3,125, 10° = +3,250, 20°
= +3,200, and 25° = +3,050.
It is evident that the decahydrated salt dissolving in water gives
a decrease of temperature. Solutions in hydrochloric acid give a
still greater decrease, because they contain the water of
crystallisation in a solid state--that is, like ice--and this on
melting absorbs heat. A mixture of 15 parts of
Na_{2}SO_{4},10H_{2}O and 12 parts of strong hydrochloric acid
produces sufficient cold to freeze water. During the treatment with
hydrochloric acid a certain quantity of sodium chloride is formed.
Sodium sulphate, Na_{2}SO_{4}, only enters into a few reactions of
combination with other salts, and chiefly with salts of the same acid,
forming double sulphates. Thus, for example, if a solution of sodium
sulphate be mixed with a solution of aluminium, magnesium, or ferrous
sulphate, it gives crystals of a double salt when evaporated. Sulphuric
acid itself forms a compound with sodium sulphate, which is exactly
like these double salts. It is formed with great ease when sodium
sulphate is dissolved in sulphuric acid and the solution evaporated.
On evaporation, crystals of the acid salt separate, Na_{2}SO_{4} +
H_{2}SO_{4} = 2NaHSO_{4}. This separates from hot solutions, whilst the
crystallo-hydrate, NaHSO_{4},H_{2}O,[10] separates from cold solutions.
The crystals when exposed to damp air decompose into H_{2}SO_{4}, which
deliquesces, and Na_{2}SO_{4} (Graham, Rose); alcohol also extracts
sulphuric acid from the acid salt. This shows the feeble force which
holds the sulphuric acid to the sodium sulphate.[11] Both acid sodium
sulphate and all mixtures of the normal salt and sulphuric acid lose
water when heated, and are converted into sodium _pyrosulphate_,
Na_{2}S_{2}O_{7}, at a low red heat.[11 bis] This anhydrous salt, at a
bright red heat, parts with the elements of sulphuric anhydride, the
normal sodium sulphate remaining behind--Na_{2}S_{2}O_{7} = Na_{2}SO_{4}
+ SO_{3}. From this it is seen that the normal salt is able to combine
with water, with other sulphates, and with sulphuric anhydride or
acid, &c.
[10] The very large and well-formed crystals of this salt resemble the
hydrate H_{2}SO_{4},H_{2}O, or SO(OH)_{4}. In general the
replacement of hydrogen by sodium modifies many of the properties
of acids less than its replacement by other metals. This most
probably depends on the volumes being nearly equal.
[11] In solution (Berthelot) the acid salt in all probability
decomposes most in the greatest mass of water. The specific
gravity (according to the determinations of Marignac) of solutions
at 15°/4° = 9,992 + 77·92_p_ + 0·231_p_^2 (_see_ Note 7). From
these figures, and from the specific gravities of sulphuric acid,
it is evident that on mixing solutions of this acid and sodium
sulphate _expansion_ will always take place; for instance,
H_{2}SO_{4} + 25H_{2}O with Na_{2}SO_{4} + 25H_{2}O increases from
483 volumes to 486. In addition to which, in weak solutions heat
is absorbed, as shown in Chapter X., Note 27. Nevertheless, even
more acid salts may be formed and obtained in a crystalline form.
For instance, on cooling a solution of 1 part of sodium sulphate
in 7 parts of sulphuric acid, crystals of the composition
NaHSO_{4},H_{2}SO_{4} are separated (Schultz, 1868). This compound
fuses at about 100°; the ordinary acid salt, NaHSO_{4}, at 149°.
[11 bis] On decreasing the pressure, sodium hydrogen sulphate,
NaHSO_{4}, dissociates much more easily than at the ordinary
pressure; it loses water and forms the pyrosulphate,
Na_{2}S_{2}O_{7}; this reaction is utilised in chemical works.
Sodium sulphate may by double decomposition be converted into a sodium
salt of any other acid, by means of heat and taking advantage of the
volatility, or by means of solution and taking advantage of the different
degree of solubility of the different salts. Thus, for instance, owing to
the insolubility of barium sulphate, sodium hydroxide or caustic soda may
be prepared from sodium sulphate, if barium hydroxide be added to its
solution, Na_{2}SO_{4} + Ba(HO)_{2} = BaSO_{4} + 2NaHO. And by taking any
salt of barium, BaX_{2}, the corresponding salt of sodium may be
obtained, Na_{2}SO_{4} + BaX_{2} = BaSO_{4} + 2NaX. Barium sulphate thus
formed, being a very sparingly-soluble salt, is obtained as a
precipitate, whilst the sodium hydroxide, or salt, NaX, is obtained in
solution, because _all salts of sodium are soluble_. Berthollet's
doctrine permits all such cases to be foreseen.
The reactions of _decomposition_ of sodium sulphate are above all
noticeable by the separation of oxygen. Sodium sulphate by itself
is very stable, and it is only at a temperature sufficient to melt
iron that it is possible to separate the elements SO_{3} from it, and
then only partially. However, the oxygen may be separated from sodium
sulphate, as from all other sulphates, by means of many substances
which are able to combine with oxygen, such as charcoal and sulphur,
but hydrogen is not able to produce this action. If sodium sulphate be
heated with charcoal, then carbonic oxide and anhydride are evolved,
and there is produced, according to the circumstances, either the
lower oxygen compound, sodium sulphite, Na_{2}SO_{3} (for instance, in
the formation of glass); or else the decomposition proceeds further,
and sodium sulphide, Na_{2}S, is formed, according to the equation
Na_{2}SO_{4} + 2C = 2CO_{2} + Na_{2}S.
On the basis of this reaction the greater part of the sulphate of sodium
prepared at chemical works is converted into _soda ash_--that is,
_sodium carbonate_, Na_{2}CO_{3}, which is used for many purposes. In
the form of carbonates, the metallic oxides behave in many cases just
as they do in the state of oxides or hydroxides, owing to the feeble
acid properties of carbonic acid. However, the majority of the salts of
carbonic acid are insoluble, whilst sodium carbonate is one of the few
soluble salts of this acid, and therefore reacts with facility. Hence
sodium carbonate is employed for many purposes, in which its alkaline
properties come into play. Thus, even under the action of feeble organic
acids it immediately parts with its carbonic acid, and gives a sodium
salt of the acid taken. Its solutions exhibit an alkaline reaction on
litmus. It aids the passage of certain organic substances (tar, acids)
into solution, and is therefore used, like caustic alkalis and soap
(which latter also acts by virtue of the alkali it contains), for the
removal of certain organic substances, especially in bleaching cotton
and similar fabrics. Besides which a considerable quantity of sodium
carbonate is used for the preparation of sodium hydroxide or caustic
soda, which has also a very wide application. In large chemical works
where sodium carbonate is manufactured from Na_{2}SO_{4}, it is usual
first to manufacture sulphuric acid, and then by its aid to convert
common salt into sodium sulphate, and lastly to convert the sodium
sulphate thus obtained into carbonate and caustic soda. Hence these
works prepare both alkaline substances (soda ash and caustic soda) and
acid substances (sulphuric and hydrochloric acids), the two classes of
chemical products which are distinguished for the greatest energy of
their reactions and are therefore most frequently applied to technical
purposes. Factories manufacturing soda are generally called alkali works.
The process of the conversion of sodium sulphate into sodium carbonate
consists in strongly heating a mixture of the sulphate with charcoal
and calcium carbonate. The following reactions then take place: the
sodium sulphate is first deoxidised by the charcoal, forming sodium
sulphide and carbonic anhydride, Na_{2}SO_{4} + 2C = Na_{2}S + 2CO_{2}.
The sodium sulphide thus formed then enters into double decomposition
with the calcium carbonate taken, and gives calcium sulphide and sodium
carbonate, Na_{2}S + CaCO_{3} = Na_{2}CO_{3} + CaS.
[Illustration: FIG. 68.--Reverberatory furnace for the manufacture
of sodium carbonate. F, grate. A, bridge. M, hearth for the ultimate
calcination of the mixture of sodium sulphate, coal, and calcium
carbonate, which is charged from above into the part of the furnace
furthest removed from the fire F. P, P, doors for stirring and bringing
the mass towards the grate F by means of stirrers R. At the end of the
operation the semifused mass is charged into trucks C.]
Besides which, under the action of the heat, a portion of the excess
of calcium carbonate is decomposed into lime and carbonic anhydride,
CaCO_{3} = CaO + CO_{2}, and the carbonic anhydride with the excess
of charcoal forms carbon monoxide, which towards the end of the
operation shows itself by the appearance of a blue flame. Thus from a
mass containing sodium sulphate we obtain a mass which includes sodium
carbonate, calcium sulphide, and calcium oxide, but none of the sodium
sulphide which was formed on first heating the mixture. The entire
process, which proceeds at a high temperature, may be expressed by a
combination of the three above-mentioned formulæ, if it be considered
that the product contains one equivalent of calcium oxide to two
equivalents of calcium sulphide.[12] The sum of the reactions may then
be expressed thus: 2Na_{2}SO_{4} + 3CaCO_{3} + 9C = 2Na_{2}CO_{3} +
CaO,2CaS + 10CO. Indeed, the quantities in which the substances are
mixed together at chemical works approaches to the proportion required
by this equation. The entire process of decomposition is carried on
in reverberatory furnaces, into which a mixture of 1,000 parts of
sodium sulphate, 1,040 parts of calcium carbonate (as a somewhat porous
limestone), and 500 parts of small coal is introduced from above.
This mixture is first heated in the portion of the furnace which is
furthest removed from the fire-grate; it is then brought to the portion
nearest to the fire-grate, when it is stirred during heating. The
partially fused mass obtained at the end of the process is cooled, and
then subjected to methodical lixiviation[13] to extract the sodium
carbonate, the mixture of calcium oxide and sulphide forming the
so-called 'soda waste' or 'alkali waste.'[14]
[12] Calcium sulphide, CaS, like many metallic sulphides which are
soluble in water, is decomposed by it (Chapter X.), CaS + H_{2}O =
CaO + H_{2}S, because hydrogen sulphide is a very feeble acid. If
calcium sulphide be acted on by a large mass of water, lime may be
precipitated, and a state of equilibrium will be reached, when the
system CaO + 2CaS remains unchanged. Lime, being a product of the
action of water on CaS, limits this action. Therefore, if in black
ash the lime were not in excess, a part of the sulphide would be
in solution (actually there is but very little). In this manner in
the manufacture of sodium carbonate the conditions of equilibrium
which enter into double decompositions have been made use of (_see
above_), and the aim is to form directly the unchangeable product
CaO,2CaS. This was first regarded as a special insoluble compound,
but there is no evidence of its independent existence.
[13] [Illustration: FIG. 69.--Apparatus for the methodical lixiviation
of black ash, &c. Water flows into the tanks from the pipes _r_,
_r_, and the saturated liquid is drawn off from _c_, _c_.]
_Methodical lixiviation_ is the extraction, by means of water, of
a soluble substance from the mass containing it. It is carried on
so as not to obtain weak aqueous solutions, and in such a way that
the residue shall not contain any of the soluble substance. This
problem is practically of great importance in many industries. It
is required to extract from the mass all that is soluble in water.
This is easily effected if water be first poured on the mass, the
strong solution thus obtained decanted, then water again poured
on, time being allowed for it to act, then again decanted, and so
on until fresh water does not take up anything. But then finally
such weak solutions are obtained that it would be very
disadvantageous to evaporate them. This is avoided by pouring the
fresh hot water destined for the lixiviation, not onto the fresh
mass, but upon a mass which has already been subjected to a first
lixiviation by weak solutions. In this way the fresh water gives a
weak solution. The strong solution which goes to the evaporating
pan flows from those parts of the apparatus which contain the
fresh, as yet unlixiviated, mass, and thus in the latter parts the
weak alkali formed in the other parts of the apparatus becomes
saturated as far as possible with the soluble substance. Generally
several intercommunicating vessels are constructed (standing at
the same level) into which in turn the fresh mass is charged which
is intended for lixiviation; the water is poured in, the alkali
drawn off, and the lixiviated residue removed. The illustration
represents such an apparatus, consisting of four communicating
vessels. The water poured into one of them flows through the two
nearest and issues from the third. The fresh mass being placed in
one of these boxes or vessels, the stream of water passing through
the apparatus is directed in such a manner as to finally issue
from this vessel containing the fresh unlixiviated mass. The fresh
water is added to the vessel containing the material which has
been almost completely exhausted. Passing through this vessel it
is conveyed by the pipe (syphon passing from the bottom of the
first box to the top of the second) communicating with the second;
it finally passes (also through a syphon pipe) into the box (the
third) containing the fresh material. The water will extract all
that is soluble in the first vessel, leaving only an insoluble
residue. This vessel is then ready to be emptied, and refilled
with fresh material. The levels of the liquids in the various
vessels will naturally be different, in consequence of the various
strengths of the solutions which they contain.
It must not, however, be thought that sodium carbonate alone
passes into the solution; there is also a good deal of caustic
soda with it, formed by the action of lime on the carbonate of
sodium, and there are also certain sodium sulphur compounds with
which we shall partly become acquainted hereafter. The sodium
carbonate, therefore, is not obtained in a very pure state. The
solution is concentrated by evaporation. This is conducted by
means of the waste heat from the soda furnaces, together with that
of the gases given off. The process in the soda furnaces can only
be carried on at a high temperature, and therefore the smoke and
gases issuing from them are necessarily very hot. If the heat they
contain was not made use of there would be a great waste of fuel;
consequently in immediate proximity to these furnaces there is
generally a series of pans or evaporating boilers, under which the
gases pass, and into which the alkali solution is poured. On
evaporating the solution, first of all the undecomposed sodium
sulphate separates, then the sodium carbonate or soda crystals.
These crystals as they separate are raked out and placed on
planks, where the liquid drains away from them. Caustic soda
remains in the residue, and also any sodium chloride which was not
decomposed in the foregoing process.
Part of the sodium carbonate is recrystallised in order to purify
it more thoroughly. In order to do this a saturated solution is
left to crystallise at a temperature below 30° in a current of
air, in order to promote the separation of the water vapour. The
large transparent crystals (efflorescent in air) of
Na_{2}CO_{3},10H_{2}O are then formed which have already been
spoken of (Chapter I.).
[14] The whole of the sulphur used in the production of the sulphuric
acid employed in decomposing the common salt is contained in this
residue. This is the great burden and expense of the soda works
which use Leblanc's method. As an instructive example from a
chemical point of view, it is worth while mentioning here two of
the various methods of recovering the sulphur from the soda waste.
Chance's process is treated in Chapter XX., Note 6.
Kynaston (1885) treats the soda waste with a solution (sp. gr.
l·21) of magnesium chloride, which disengages sulphuretted
hydrogen: CaS + MgCl_{2} + 2H_{2}O = CaCl_{2} + Mg(OH)_{2} +
H_{2}S. Sulphurous anhydride is passed through the residue in
order to form the insoluble calcium sulphite: CaCl_{2} +
Mg(OH)_{2} + SO_{2} = CaSO_{3} + MgCl_{2} + H_{2}O. The solution
of magnesium chloride obtained is again used, and the washed
calcium sulphite is brought into contact at a low temperature with
hydrochloric acid (a weak aqueous solution) and hydrogen sulphide,
the whole of the sulphur then separating:
CaSO_{3} + 2H_{2}S + 2HCl = CaCl_{2} + 3H_{2}O + 3S.
But most efforts have been directed towards avoiding the formation
of soda waste.
The above-mentioned process for making soda was discovered in the year
1808 by the French doctor Leblanc, and is known as the Leblanc process.
The particulars of the discovery are somewhat remarkable. Sodium
carbonate, having a considerable application in industry, was for a
long time prepared exclusively from the ash of marine plants (Chapter
XI., page 497). Even up to the present time this process is carried on
in Normandy. In France, where for a long time the manufacture of large
quantities of soap (so-called Marseilles soap) and various fabrics
required a large amount of soda, the quantity prepared at the coast was
insufficient to meet the demand. For this reason during the wars at the
beginning of the century, when the import of foreign goods into France
was interdicted, the want of sodium carbonate was felt. The French
Academy offered a prize for the discovery of a profitable method of
preparing it from common salt. Leblanc then proposed the above-mentioned
process, which is remarkable for its great simplicity.[15]
[15] Among the drawbacks of the Leblanc process are the accumulation of
'soda waste' (Note 14) owing to the impossibility at the
comparatively low price of sulphur (especially in the form of
pyrites) of finding employment for the sulphur and sulphur
compounds for which this waste is sometimes treated, and also the
insufficient purity of the sodium carbonate for many purposes. The
advantages of the Leblanc process, besides its simplicity and
cheapness, are that almost the whole of the acids obtained as
bye-products have a commercial value; for chlorine and bleaching
powder are produced from the large amount of hydrochloric acid
which appears as a bye-product; caustic soda also is very easily
made, and the demand for it increases every year. In those places
where salt, pyrites, charcoal, and limestone (the materials
required for alkali works) are found side by side--as, for
instance, in the Ural or Don districts--conditions are favourable
to the development of the manufacture of sodium carbonate on an
enormous scale; and where, as in the Caucasus, sodium sulphate
occurs naturally, the conditions are still more favourable. A
large amount, however, of the latter salt, even from soda works,
is used in making glass. The most important soda works, as regards
the quantity of products obtained from them, are the English
works.
As an example of the other numerous and varied methods of
manufacturing soda from sodium chloride, the following may be
mentioned: Sodium chloride is decomposed by oxide of lead, PbO,
forming lead chloride and sodium oxide, which, with carbonic
anhydride, yields sodium carbonate (Scheele's process). In Cornu's
method sodium chloride is treated with lime, and then exposed to
the air, when it yields a small quantity of sodium carbonate. In
E. Kopp's process sodium sulphate (125 parts) is mixed with oxide
of iron (80 parts) and charcoal (55 parts), and the mixture is
heated in reverberatory furnaces. Here a compound,
Na_{6}Fe_{4}S_{3}, is formed, which is insoluble in water absorbs
oxygen and carbonic anhydride, and then forms sodium carbonate and
ferrous sulphide; this when roasted gives sulphurous anhydride,
the indispensable material for the manufacture of sulphuric acid,
and ferric oxide which is again used in the process. In Grant's
method sodium sulphate is transformed into sodium sulphide, and
the latter is decomposed by a stream of carbonic anhydride and
steam, when hydrogen sulphide is disengaged and sodium carbonate
formed. Gossage prepares Na_{2}S from Na_{2}SO_{4} (by heating it
with carbon), dissolves it in water and subjects the solution to
the action of an excess of CO_{2} in coke towers, thus obtaining
H_{2}S (a gas which gives SO_{2} under perfect combustion, or
sulphur when incompletely burnt, Chapter XX., Note 6) and
bicarbonate of sodium; Na_{2}S + 2CO_{2} + 2H_{2}O = H_{2}S +
2HNaCO_{3}. The latter gives soda and CO_{2} when ignited. This
process quite eliminates the formation of soda-waste (_see_ Note
3) and should in my opinion be suitable for the treatment of
native Na_{2}SO_{4}, like that which is found in the Caucasus, all
the more since H_{2}S gives sulphur as a bye-product.
Repeated efforts have been made in recent times to obtain soda
(and chlorine, _see_ Chapter II., Note 1) from strong solutions of
salt (Chapter X., Note 23 bis) by the action of an electric
current, but until now these methods have not been worked out
sufficiently for practical use, probably partly owing to the
complicated apparatus needed, and the fact that the chlorine given
off at the anode corrodes the electrodes and vessels and has but a
limited industrial application. We may mention that according to
Hempel (1890) soda in crystals is deposited when an electric
current and a stream of carbonic acid gas are passed through a
saturated solution of NaCl.
Sodium carbonate may likewise be obtained from cryolite (Chapter
XVII., Note 23) the method of treating this will be mentioned
under Aluminium.
Of all other industrial processes for manufacturing sodium carbonate,
the _ammonia process_ is the most worthy of mention.[16] In this the
vapours of ammonia, and then an excess of carbonic anhydride, are
directly introduced into a concentrated solution of sodium chloride
in order to form the acid ammonium carbonate, NH_{4}HCO_{3}. Then, by
means of the double saline decomposition of this salt, sodium chloride
is decomposed, and in virtue of its slight solubility acid sodium
carbonate, NaHCO_{3}, is precipitated and ammonium chloride, NH_{4}Cl,
is obtained in solution (with a portion of the sodium chloride and acid
sodium carbonate). The reaction proceeds in the solution owing to the
sparing solubility of the NaHCO_{3} according to the equation NaCl +
NH_{4}HCO_{3} = NH_{4}Cl + NaHCO_{3}. The ammonia is recovered from the
solution by heating with lime or magnesia,[16 bis] and the precipitated
acid sodium carbonate is converted into the normal salt by heating. It
is thus obtained in a very pure state.[17]
[16] This process (Chapter XVII.) was first pointed out by Turck,
worked out by Schloesing, and finally applied industrially by
Solvay. The first (1883) large soda factories erected in Russia
for working this process are on the banks of the Kama at
Berezniak, near Ousolia, and belong to Lubimoff. But Russia, which
still imports from abroad a large quantity of bleaching powder and
exports a large amount of manganese ore, most of all requires
works carrying on the Leblanc process. In 1890 a factory of this
kind was erected by P. K. Oushkoff, on the Kama, near Elagoubi.
[16 bis] Mond (_see_ Chapter XI., Note 3 bis) separates the NH_{4}Cl
from the residual solutions by cooling (Chapter X., Note 44);
ignites the sal-ammoniac and passes the vapour over MgO, and so
re-obtains the NH_{3}, and forms MgCl_{2}: the former goes back
for the manufacture of soda, while the latter is employed either
for making HCl or Cl_{2}.
[17] Commercial soda ash (calcined, anhydrous) is rarely pure; the
crystallised soda is generally purer. In order to purify it
further, it is best to boil a concentrated solution of soda ash
until two-thirds of the liquid remain, collect the soda which
settles, wash with cold water, and then shake up with a strong
solution of ammonia, pour off the residue, and heat. The
impurities will then remain in the mother liquors, &c.
Some numerical data may be given for sodium carbonate. The
specific gravity of the anhydrous salt is 2·48, that of the
decahydrated salt 1·46. Two varieties are known of the
heptahydrated salt (Löwel, Marignac, Rammelsberg), which are
formed together by allowing a saturated solution to cool under a
layer of alcohol; the one is less stable (like the corresponding
sulphate) and at 0° has a solubility of 32 parts (of anhydrous
salt) in 100 water; the other is more stable, and its solubility
20 parts (of anhydrous salt) per 100 of water. The solubility of
the decahydrated salt in 100 water = at 0°, 7·0; at 20°, 21·7; at
30°, 37·2 parts (of anhydrous salt). At 80° the solubility is only
46·1, at 90° 45·7, at 100°, 45·4 parts (of anhydrous salt). That
is, it falls as the temperature rises, like Na_{2}SO_{4}. The
specific gravity (Note 7) of the solutions of sodium carbonate,
according to the data of Gerlach and Kohlrausch, at 15°/4° is
expressed by the formula, _s_ = 9,992 + 104·5_p_ + 0·165_p_^2.
Weak solutions occupy a volume not only less than the sum of the
volumes of the anhydrous salt and the water, but even less than
the water contained in them. For instance, 1,000 grams of a 1 p.c.
solution occupy (at 15°) a volume of 990·4 c.c. (sp. gr. 1·0097),
but contain 990 grams of water, occupying at 15° a volume of 990·8
c.c. A similar case, which is comparatively rare occurs also with
sodium hydroxide, in those dilute solutions for which the factor
_A_ is greater than 100 if the sp. gr. of water at 4° = 100,000,
and if the sp. gr. of the solution be expressed by the formula _S_
= _S__{0} + _Ap_ + _Bp_^2, where _S__{0} is the specific gravity
of the water. For 5 p.c. the sp. gr. 15°/4° = 1·0520; for 10 p.c.
1·1057; for 15 p.c. 1·1603. The changes in the sp. gr. with the
temperature are here almost the same as with solutions of sodium
chloride with an equal value of _p_.
Sodium carbonate, like sodium sulphate, loses all its water on being
heated, and when anhydrous fuses at a bright-red heat (1098°). A small
quantity of sodium carbonate placed in the loop of a platinum wire
volatilises in the heat of a gas flame, and therefore in the furnaces of
glass works part of the soda is always transformed into the condition
of vapour. Sodium carbonate resembles sodium sulphate in its relation
to water.[18] Here also the greatest solubility is at the temperature
of 37°; both salts, on crystallising at the ordinary temperature,
combine with ten molecules of water, and such crystals of soda, like
crystals of Glauber's salt, fuse at 34°. Sodium carbonate also forms a
supersaturated solution, and, according to the conditions, gives various
combinations with water of crystallisation (mentioned on page 108), &c.
[18] The resemblance is so great that, notwithstanding the difference
in the molecular composition of Na_{2}SO_{4} and Na_{2}CO_{3},
they ought to be classed under the type (NaO)_{2}R, where R =
SO_{2} or CO. Many other sodium salts also contain 10 mol. H_{2}O.
At a red heat superheated steam liberates carbonic anhydride from
sodium carbonate and forms caustic soda, Na_{2}CO_{3} + H_{2}O = 2NaHO
+ CO_{2}. Here the carbonic anhydride is replaced by water; this
depends on the feebly acid character of carbonic anhydride. By direct
heating, sodium carbonate is only slightly decomposed into sodium oxide
and carbonic anhydride; thus, when sodium carbonate is fused, about 1
per cent. of carbonic anhydride is disengaged.[19] The carbonates of
many other metals--for instance, of calcium, copper, magnesium, iron,
&c.--on being heated lose all their carbonic anhydride. This shows the
considerable basic energy which sodium possesses. With the soluble salts
of most metals, sodium carbonate gives precipitates either of insoluble
carbonates of the metals, or else of the hydroxides (in this latter
case carbonic anhydride is disengaged); for instance, with barium salts
it precipitates an insoluble barium carbonate (BaCl_{2} + Na_{2}CO_{3}
= 2NaCl + BaCO_{3}) and with the aluminium salts it precipitates
aluminium hydroxide, carbonic anhydride being disengaged: 3Na_{2}CO_{3}
+ Al_{2}(SO_{4})_{3} + 3H_{2}O = 3Na_{2}SO_{4} + 2Al(OH)_{3} + 3CO_{2}.
Sodium carbonate, like all the salts of carbonic acid, evolves carbonic
anhydride on treatment with all acids which are to any extent energetic.
But if an acid diluted with water be gradually added to a solution of
sodium carbonate, _at first_ such an evolution does not take place,
because the excess of the carbonic anhydride forms acid sodium carbonate
(sodium bicarbonate), NaHCO_{3}.[20] The acid sodium carbonate is an
unstable salt. Not only when heated alone, but even on being slightly
heated in solution, and also at the ordinary temperature in damp air,
it loses carbonic anhydride and forms the normal salt. And at the
same time it is easy to obtain it in a pure crystalline form, if a
strong solution of sodium carbonate be cooled and a stream of carbonic
anhydride gas passed through it. The acid salt is less soluble in water
than the normal,[21] and therefore a strong solution of the latter
gives crystals of the acid salt if carbonic anhydride be passed through
it. The acid salt may be yet more conveniently formed from effloresced
crystals of sodium carbonate, which, on being considerably heated, very
easily absorb carbonic anhydride.[22] The acid salt crystallises well,
but not, however, in such large crystals as the normal salt; it has a
brackish and not an alkaline taste like that of the normal salt; its
reaction is feebly alkaline, nearly neutral. At 70° its solution begins
to lose carbonic anhydride, and on boiling the evolution becomes very
abundant. From the preceding remarks it is clear that in most reactions
this salt, especially when heated, acts similarly to the normal salt,
but has, naturally, some distinction from it. Thus, for example, if a
solution of sodium carbonate be added to a normal magnesium salt, a
turbidity (precipitate) is formed of magnesium carbonate. MgCO_{3}. No
such precipitate is formed by the acid salt, because magnesium carbonate
is soluble in the presence of an excess of carbonic anhydride.
[19] According to the observations of Pickering. According to Rose,
when solutions of sodium carbonate are boiled a certain amount of
carbonic anhydride is disengaged.
[20] The composition of this salt, however, may be also represented
as a combination of carbonic acid, H_{2}CO_{3}, with the normal
salt, Na_{2}CO_{3}, just as the latter also combines with water.
Such a combination is all the more likely because (1) there exists
another salt, Na_{2}CO_{3},2NaHCO_{3},2H_{2}O (sodium
sesquicarbonate), obtained by cooling a boiling solution of sodium
bicarbonate, or by mixing this salt with the normal salt; but the
formula of this salt cannot be derived from that of normal
carbonic acid, as the formula of the bicarbonate can. At the same
time the sesqui-salt has all the properties of a definite
compound; it crystallises in transparent crystals, has a constant
composition, its solubility (at 0° in 100 of water, 12·6 of
anhydrous salt) differs from the solubility of the normal and acid
salts; it is found in nature, and is known by the names of _trona_
and _urao_. The observations of Watts and Richards showed (1886)
that on pouring a strong solution of the acid salt into a solution
of the normal salt saturated by heating, crystals of the salt
NaHCO_{3},Na_{2}CO_{3},2H_{2}O may be easily obtained, as long as
the temperature is above 35°. The natural urao (Boussingault) has,
according to Laurent, the same composition. This salt is very
stable in air, and may be used for purifying sodium carbonate on
the large scale. Such compounds have been little studied from a
theoretical point of view, although particularly interesting,
since in all probability they correspond with ortho-carbonic acid,
C(OH)_{4}, and at the same time correspond with double salts like
astrakhanite (Chapter XIV., Note 25). (2) Water of crystallisation
does not enter into the composition of the crystals of the acid
salt, so that on its formation (occurring only at low
temperatures, as in the formation of crystalline compounds with
water) the water of crystallisation of the normal salt separates
and the water is, as it were, replaced by the elements of carbonic
acid. If anhydrous sodium carbonate be mixed with the amount of
water requisite for the formation of Na_{2}CO_{3},H_{2}O, this
salt will, when powdered, absorb CO_{2} as easily at the ordinary
temperature as it does water.
[21] 100 parts of water at 0° dissolve 7 parts of the acid salt, which
corresponds with 4·3 parts of the anhydrous normal salt, but at 0°
100 parts of water dissolve 7 parts of the latter. The solubility
of the bi-or acid salt varies with considerable regularity; 100
parts of water dissolves at 15° 9 parts of the salt, at 30° 11
parts.
The ammonium, and more especially the calcium, salt, is much more
soluble in water. The ammonia process (_see_ p. 524) is founded
upon this. Ammonium bicarbonate (acid carbonate) at 0° has a
solubility of 12 parts in 100 water, at 30° of 27 parts. The
solubility therefore increases very rapidly with the temperature.
And its saturated solution is more stable than a solution of
sodium bicarbonate. In fact, saturated solutions of these salts
have a gaseous tension like that of a mixture of carbonic
anhydride and water--namely, at 15° and at 50°, for the sodium
salt 120 and 750 millimetres, for the ammonium salt 120 and 563
millimetres. These data are of great importance in understanding
the phenomena connected with the ammonia process. They indicate
that with an increased pressure the formation of the sodium salt
ought to increase if there be an excess of ammonium salt.
[22] Crystalline sodium carbonate (broken into lumps) also absorbs
carbonic anhydride, but the water contained in the crystals is
then disengaged: Na_{2}CO_{3},10H_{2}O + CO_{2} =
Na_{2}CO_{3},H_{2}CO_{3} + 9H_{2}O, and dissolves part of the
carbonate; therefore part of the sodium carbonate passes into
solution together with all the impurities. When it is required to
avoid the formation of this solution, a mixture of ignited and
crystalline sodium carbonate is taken. Sodium bicarbonate is
prepared chiefly for medicinal use, and is then often termed
_carbonate of soda_, also, for instance, in the so-called soda
powders, for preparing certain artificial mineral waters, for the
manufacture of digestive lozenges like those made at Essentuki,
Vichy, &c.
Sodium carbonate is used for the preparation of _caustic soda_[23]--that
is, the hydrate of sodium oxide, or the alkali which corresponds to
sodium. For this purpose the action of lime on a solution of sodium
carbonate is generally made use of. The process is as follows: a weak,
generally 10 per cent., solution of sodium carbonate is taken,[24] and
boiled in a cast-iron, wrought-iron, or silver boiler (sodium hydroxide
does not act on these metals), and lime is added, little by little,
during the boiling. This latter is soluble in water, although but very
slightly. The clear solution becomes turbid on the addition of the lime
because a precipitate is formed; this precipitate consists of calcium
carbonate, almost insoluble in water, whilst caustic soda is formed
and remains in solution. The decomposition is effected according to
the equation: Na_{2}CO_{3} + Ca(HO)_{2} = CaCO_{3} + 2NaHO. On cooling
the solution the calcium carbonate easily settles as a precipitate,
and the clear solution or alkali above it contains the easily soluble
sodium hydroxide formed in the reaction.[25] After the necessary
quantity of lime has been added, the solution is allowed to stand, and
is then decanted off and evaporated in cast or wrought iron boilers,
or in silver pans if a perfectly pure product is required.[26] The
evaporation cannot be conducted in china, glass, or similar vessels,
because caustic soda attacks these materials, although but slightly. The
solution does not crystallise on evaporation, because the solubility
of caustic soda when hot is very great, but crystals containing water
of crystallisation may be obtained by cooling. If the evaporation
of the alkali be conducted until the specific gravity reaches 1·38,
and the liquid is then cooled to 0°, transparent crystals appear
containing 2NaHO,7H_{2}O; they fuse at +6°.[27] If the evaporation be
conducted so long as water is disengaged, which requires a considerable
amount of heat, then, on cooling, the hydroxide, NaHO, solidifies
in a semi-transparent crystalline mass,[28] which eagerly absorbs
moisture and carbonic anhydride from the air.[29] Its specific gravity
is 2·13;[30] it is easily soluble in water, with disengagement of a
considerable quantity of heat.[31] A saturated solution at the ordinary
temperature has a specific gravity of about 1·5, contains about 45 per
cent. of sodium hydroxide, and boils at 130°; at 55° water dissolves
an equal weight of it.[32] Caustic soda is not only soluble in water
but in alcohol, and even in ether. Dilute solutions of sodium hydroxide
produce a soapy feeling on the skin because the active base of soap
consists of caustic soda.[33] Strong solutions have a corroding action.
[23] In chemistry, sodium oxide is termed 'soda,' which word must be
carefully distinguished from the word sodium, meaning the metal.
[24] With a small quantity of water, the reaction either does not
take place, or even proceeds in the reverse way--that is, sodium
and potassium hydroxides remove carbonic anhydride from calcium
carbonate (Liebig, Watson, Mitscherlich, and others). The
influence of the mass of water is evident. According to Gerberts,
however, strong solutions of sodium carbonate are decomposed by
lime, which is very interesting if confirmed by further
investigation.
[25] As long as any undecomposed sodium carbonate remains in solution,
excess of acid added to the solution disengages carbonic
anhydride, and the solution after dilution gives a white
precipitate with a barium salt soluble in acids, showing the
presence of a carbonate in solution (if there be sulphate present,
it also forms a white precipitate, but this is insoluble in
acids). For the decomposition of sodium carbonate, milk of
lime--that is, slaked slime suspended in water--is employed.
Formerly pure sodium hydroxide was prepared (according to
Berthollet's process) by dissolving the impure substance in
alcohol (sodium carbonate and sulphate are not soluble), but now
that metallic sodium has become cheap and is purified by
distillation, _pure caustic soda_ is prepared by acting on a small
quantity of water with sodium. Perfectly pure sodium hydroxide may
also be obtained by allowing strong solutions to crystallise (in
the cold) (Note 27).
In alkali works where the Leblanc process is used, caustic soda is
prepared directly from the alkali remaining in the mother liquors
after the separation of the sodium carbonate by evaporation (Note
14). If excess of lime and charcoal have been used, much sodium
hydroxide maybe obtained. After the removal as much as possible of
the sodium carbonate, a red liquid (from iron oxide) is left,
containing sodium hydroxide mixed with compounds of sulphur and of
cyanogen (_see_ Chapter IX.) and also containing iron. This red
alkali is evaporated and air is blown through it, which oxidises
the impurities (for this purpose sometimes sodium nitrate is
added, or bleaching powder, &c.) and leaves fused caustic soda.
The fused mass is allowed to settle in order to separate the
ferruginous precipitate, and poured into iron drums, where the
sodium hydroxide solidifies. Such caustic soda contains about 10
p.c. of water in excess and some saline impurities, but when
properly manufactured is almost free from carbonate and from iron.
The greater part of the caustic soda, which forms so important an
article of commerce, is manufactured in this manner.
[26] Löwig gave a method of preparing sodium hydroxide from sodium
carbonate by heating it to a dull red heat with an excess of
ferric oxide. Carbonic anhydride is given off, and warm water
extracts the caustic soda from the remaining mass. This reaction,
as experiment shows, proceeds very easily, and is an example of
contact action similar to that of ferric oxide on the
decomposition of potassium chlorate. The reason of this may be
that a small quantity of the sodium carbonate enters into double
decomposition with the ferric oxide, and the ferric carbonate
produced is decomposed into carbonic anhydride and ferric oxide,
the action of which is renewed. Similar explanations expressing
the _reason_ for a reaction really adds but little to that
elementary conception of contact which, according to my opinion,
consists in the change of motion of the atoms in the molecules
under the influence of the substance in contact. In order to
represent this clearly it is sufficient, for instance, to imagine
that in the sodium carbonate the elements CO_{2} move in a circle
round the elements Na_{2}O, but at the points of contact with
Fe_{2}O_{3} the motion becomes elliptic with a long axis, and at
some distance from Na_{2}O the elements of CO_{2} are parted, not
having the faculty of attaching themselves to Fe_{2}O_{3}.
[27] By allowing strong solutions of sodium hydroxide to crystallise in
the cold, impurities--such as, for instance, sodium sulphate--may
be separated from them. The fused crystallo-hydrate 2NaHO,7H_{2}O
forms a solution having a specific gravity of 1·405 (Hermes). The
crystals on dissolving in water produce cold, while NaHO produces
heat. Besides which Pickering obtained hydrates with 1, 2, 4, 5,
and 7 H_{2}O.
[28] In solid caustic soda there is generally an excess of water beyond
that required by the formula NaHO. The caustic soda used in
laboratories is generally cast in sticks, which are broken into
pieces. It must be preserved in carefully closed vessels, because
it absorbs water and carbonic anhydride from the air.
[29] By the way it changes in air it is easy to distinguish caustic
soda from caustic potash, which in general resembles it. Both
alkalis absorb water and carbonic anhydride from the air, but
caustic potash forms a deliquescent mass of potassium carbonate,
whilst caustic soda forms a dry powder of efflorescent salt.
[30] As the molecular weight of NaHO = 40, the volume of its molecule =
40/2·13 = 18·5, which very nearly approaches the volume of a
molecule of water. The same rule applies to the compounds of
sodium in general--for instance, its salts have a molecular volume
approaching the volume of the acids from which they are derived.
[31] The molecular quantity of sodium hydroxide (40 grams), on being
dissolved in a large mass (200 gram molecules) of water, develops,
according to Berthelot 9,780, and according to Thomsen 9,940,
heat-units, but at 100° about 13,000 (Berthelot). Solutions of
NaHO + _n_H_{2}O, on being mixed with water, evolve heat if they
contain less than 6H_{2}O, but if more they absorb beat.
[32] The specific gravity of solutions of sodium hydroxide at 15°/4° is
given in the short table below:--
NaHO, p.c. 5 10 15 20 30 40
Sp. gr. 1·057 1·113 1·169 1·224 1·331 1·436
1,000 grams of a 5 p.c. solution occupies a volume of 946 c.c.;
that is, less than the water serving to make the solution (_see_
Note 18).
[33] Sodium hydroxide and some other alkalis are capable of
hydrolysing--saponifying, as it is termed--the compounds of acids
with alcohols. If RHO (or R(HO)_{_n_}) represent the composition
of an alcohol--that is, of the hydroxide of a hydrocarbon
radicle--and QHO an acid, then the compound of the acid with the
alcohol or ethereal salt of the given acid will have the
composition RQO. Ethereal salts, therefore, present a likeness to
metallic salts, just as alcohols resemble basic hydroxides. Sodium
hydroxide acts on ethereal salts in the same way that it acts on
the majority of metallic salts--namely, it liberates alcohol, and
forms the sodium salt of that acid which was in the ethereal salt.
The reaction takes place in the following way:--
RQO + NaHO = NaQO + RHO
Ethereal Caustic Sodium Alcohol
salt soda salt
Such a decomposition is termed saponification; similar reactions
were known very long ago for the ethereal salts corresponding with
glycerin, C_{3}H_{5}(OH)_{3} (Chapter IX.), found in animals and
plants, and composing what are called fats or oils. Caustic soda,
acting on fat and oil, forms glycerin, and sodium salts of those
acids which were in union with the glycerin in the fat, as
Chevreul showed at the beginning of this century. The sodium salts
of the fatty acids are commonly known as soaps. That is to say,
soap is made from fat and caustic soda, glycerin being separated
and a sodium salt or soap formed. As glycerin is usually found in
union with certain acids, so also are the sodium salts of the same
acids found in soap. The greater part of the acids found in
conjunction with glycerin in fats are the solid palmitic and
stearic acids, C_{16}H_{32}O_{2} and C_{18}H_{38}O_{2}, and the
liquid oleic acid, C_{18}H_{34}O_{2}. In preparing soap the fatty
substances are mixed with a solution of caustic soda until an
emulsion is formed; the proper quantity of caustic soda is then
added in order to produce saponification on heating, the soap
being separated from the solution either by means of an excess of
caustic soda or else by common salt, which displaces the soap from
the aqueous solution (salt water does not dissolve soap, neither
does it form a lather). Water acting on soap partly decomposes it
(because the acids of the soap are feeble), and the alkali set
free acts during the application of soap. Hence it may be replaced
by a very feeble alkali. Strong solutions of alkali corrode the
skin and tissues. They are not formed from soap, because the
reaction is reversible, and the alkali is only set free by the
excess of water. Thus we see how the teaching of Berthollet
renders it possible to understand many phenomena which occur in
every-day experience (_see_ Chapter IX., Note 15).
The chemical _reactions of sodium hydroxide_ serve as a type for those
of a whole class of alkalis--that is, of soluble basic hydroxides,
MOH. The solution of sodium hydroxide is a very caustic liquid--that
is to say, it acts in a destructive way on most substances, for
instance on most organic tissues--hence caustic soda, like all soluble
alkalis, is a poisonous substance; acids, for example hydrochloric,
serve as antidotes. The action of caustic soda on bones, fat, starch,
and similar vegetable and animal substances explains its action on
organisms. Thus bones, when plunged into a weak solution of caustic
soda, fall to powder,[34] and evolve a smell of ammonia, owing to the
caustic soda changing the gelatinous organic substance of the bones
(which contains carbon, hydrogen, nitrogen, oxygen, and sulphur, like
albumin), dissolving it and in part destroying it, whence ammonia is
disengaged. Fats, tallow, and oils become saponified by a solution of
caustic soda--that is to say, they form with it _soaps_ soluble in
water, or sodium salts of the organic acids contained in the fats.[35]
The most characteristic reactions of sodium hydroxide are determined by
the fact that it _saturates all acids, forming salts with them_, which
are almost all soluble in water, and in this respect caustic soda is as
characteristic amongst the bases as nitric acid is among the acids. It
is impossible to detect sodium by means of the formation of precipitates
of insoluble sodium salts, as may be done with other metals, many of
whose salts are but slightly soluble. The powerful alkaline properties
of caustic soda determine its capacity for combining with even the
feeblest acids, its property of disengaging ammonia from ammonium salts,
its faculty of forming precipitates from solutions of salts whose bases
are insoluble in water, &c. If a solution of the salt of almost any
metal be mixed with caustic soda, then a soluble sodium salt will be
formed, and an insoluble hydroxide of the metal will be separated--for
instance, copper nitrate yields copper hydroxide, Cu(NO_{3})_{2} + 2NaHO
= Cu(HO)_{2} + 2NaNO_{3}. Even many _basic oxides_ precipitated by
caustic soda _are capable_ of _combining_ with it and forming soluble
compounds, and therefore caustic soda in the presence of salts of such
metals first forms a precipitate of hydroxide, and then, employed in
excess, dissolves this precipitate. This phenomenon occurs, for example,
when caustic soda is added to the salts of aluminium. This shows the
property of such an alkali as caustic soda of combining not only with
acids, but also with feeble basic oxides. For this reason caustic soda
_acts on most elements_ which are capable of forming acids or oxides
similar to them; thus the metal aluminium gives hydrogen with caustic
soda in consequence of the formation of alumina, which combines with
the caustic soda--that is, in this case, the caustic alkali acts on the
metal just as sulphuric acid does on Fe or Zn. If caustic soda acts
in this manner on a metalloid capable of combining with the hydrogen
evolved (aluminium does not give a compound with hydrogen), then it
forms such a hydrogen compound. Thus, for instance, phosphorus acts in
this way on caustic soda, yielding hydrogen phosphide. When the hydrogen
compound disengaged is capable of combining with the alkali, then,
naturally, a salt of the corresponding acid is formed. For example,
chlorine and sulphur act in this way on caustic soda. Chlorine, with the
hydrogen of the caustic soda, forms hydrochloric acid, and the latter
forms common salt with the sodium hydroxide, whilst the other atom in
the molecule of chlorine, Cl_{2}, takes the place of the hydrogen, and
forms the hypochlorite, NaClO. In the same way, by the action of sodium
hydroxide on sulphur, hydrogen sulphide is formed, which acts on the
soda forming sodium _sulphide_, in addition to which sodium thiosulphate
is formed (_see_ Chapter XX.) By virtue of such reactions, sodium
hydroxide acts on many metals and non-metals. Such action is often
accelerated by the presence of the oxygen of the air, as by this means
the formation of acids and oxides rich in oxygen is facilitated. Thus
many metals and their lower oxides, in the presence of an alkali, absorb
oxygen and form acids. Even manganese peroxide, when mixed with caustic
soda, is capable of absorbing the oxygen of the air, and forming sodium
manganate. Organic acids when heated with caustic soda give up to it the
elements of carbonic anhydride, forming sodium carbonate, and separating
that hydrocarbon group which exists, in combination with carbonic
anhydride, in the organic acid.
[34] On this is founded the process of Henkoff and Engelhardt for
treating bones. The bones are mixed with ashes, lime, and water;
it is true that in this case more potassium hydroxide than sodium
hydroxide is formed, but their action is almost identical.
[35] As explained in Note 33.
Thus sodium hydroxide, like the soluble alkalis in general, ranks
amongst the most active substances in the chemical sense of the term,
and but few substances are capable of resisting it. Even siliceous
rocks, as we shall see further on, are transformed by it, forming when
fused with it vitreous slags. Sodium hydroxide (like ammonium and
potassium hydroxides), as a typical example of the basic hydrates, in
distinction from many other basic oxides, easily _forms acid salts_
with acids (for instance, NaHSO_{4}, NaHCO_{3}), and does not form any
basic salts at all; whilst many less energetic bases, such as the oxides
of copper and lead, easily form basic salts, but acid salts only with
difficulty. This capability of forming acid salts, particularly with
polybasic acids, may be explained by the energetic basic properties
of sodium hydroxide, contrasted with the small development of these
properties in the bases which easily form basic salts. An energetic
base is capable of retaining a considerable quantity of acid, which a
slightly energetic base would not have the power of doing. Also, as will
be shown in the subsequent chapters, sodium belongs to the univalent
metals, being exchangeable for hydrogen atom for atom--that is, amongst
metals sodium may, like chlorine amongst the non-metals, serve as the
representative of the univalent properties. Most of the elements which
are not capable of forming acid salts are bivalent. Whence it may be
understood that in a bibasic acid--for instance, carbonic, H_{2}CO_{3},
or sulphuric, H_{2}SO_{4}--the hydrogen may be exchanged, atom for atom,
for sodium, and yield an acid salt by means of the first substitution,
and a normal salt by means of the second--for instance, NaHSO_{4}, and
Na_{2}SO_{4}, whilst such bivalent metals as calcium and barium do not
form acid salts because one of their atoms at once takes the place of
both hydrogen atoms, forming, for example, CaCO_{3} and CaSO_{4}.[35 bis]
[35 bis] It might be expected, from what has been mentioned above, that
bivalent metals would easily form acid salts with acids containing
more than two atoms of hydrogen--for instance, with tribasic
acids, such as phosphoric acid, H_{3}PO_{4}--and actually such
salts do exist; but all such relations are complicated by the fact
that the character of the base very often changes and becomes
weakened with the increase of valency and the change of atomic
weight; the feebler bases (like silver oxide), although
corresponding with univalent metals, do not form acid salts, while
the feeblest bases (CuO, PbO, &c.) easily form basic salts, and
notwithstanding their valency do not form acid salts which are in
any degree stable--that is, which are undecomposable by water.
Basic and acid salts ought to be regarded rather as compounds
similar to crystallo-hydrates, because such acids as sulphuric
form with sodium not only an acid and a normal salt, as might be
expected from the valency of sodium, but also salts containing a
greater quantity of acid. In sodium sesquicarbonate we saw an
example of such compounds. Taking all this into consideration, we
must say that the property of more or less easily forming acid
salts depends more upon the energy of the base than upon its
valency, and the best statement is that _the capacity of a base
for forming acid and basic salts is characteristic_, just as the
faculty of forming compounds with hydrogen is characteristic of
elements.
We have seen the transformation of common salt into sodium sulphate,
of this latter into sodium carbonate, and of sodium carbonate into
caustic soda. Lavoisier still regarded sodium hydroxide as an element,
because he was unacquainted with its decomposition with the formation
of metallic sodium, which separates the hydrogen from water, reforming
caustic soda.
The preparation of _metallic sodium_ was one of the greatest discoveries
in chemistry, not only because through it the conception of elements
became broader and more correct, but especially because in sodium,
chemical properties were observed which were but feebly shown in the
other metals more familiarly known. This discovery was made in 1807
by the English chemist _Davy_ by means of the galvanic current. By
connecting with the positive pole (of copper or carbon) a piece of
caustic soda (moistened in order to obtain electrical conductivity),
and boring a hole in it filled with mercury connected with the negative
pole of a strong Volta's pile, Davy observed that on passing the current
a peculiar metal dissolved in the mercury, less volatile than mercury,
and capable of decomposing water, again forming caustic soda. In this
way (by analysis and synthesis) Davy demonstrated the compound nature
of alkalis. On being decomposed by the galvanic current, caustic soda
disengages hydrogen and sodium at the negative pole and oxygen at the
positive pole. Davy showed that the metal formed volatilises at a red
heat, and this is its most important physical property in relation to
its extraction, all later methods being founded on it. Besides this Davy
observed that sodium easily oxidises, its vapour taking fire in air,
and the latter circumstance was for a long time an obstacle to the easy
preparation of this metal. The properties of sodium were subsequently
more thoroughly investigated by Gay-Lussac and Thénard, who observed
that metallic iron at a high temperature was capable of reducing caustic
soda to sodium.[36] Brunner latterly discovered that not only iron,
but also charcoal, has this property, although hydrogen has not.[37]
But still the methods of extracting sodium were very troublesome,
and consequently it was a great rarity. The principal obstacle
to its production was that an endeavour was made to condense the
easily-oxidising vapours of sodium in vacuo in complicated apparatus.
For this reason, when Donny and Maresca, having thoroughly studied the
matter, constructed a specially simple condenser, the production of
sodium was much facilitated. Furthermore, in practice the most important
epoch in the history of the production of sodium is comprised in the
investigation of Sainte-Claire Deville, who avoided the complex methods
in vogue up to that time, and furnished those simple means by which the
production of sodium is now rendered feasible in chemical works.
[36] Deville supposes that such a decomposition of sodium hydroxide
by metallic iron depends solely on the dissociation of the alkali
at a white heat into sodium, hydrogen, and oxygen. Here the part
played by the iron is only that it retains the oxygen formed,
otherwise the decomposed elements would again reunite upon
cooling, as in other cases of dissociation. If it be supposed that
the temperature at the commencement of the dissociation of the
iron oxides is higher than that of sodium oxide, then the
decomposition may be explained by Deville's hypothesis. Deville
demonstrates his views by the following experiment:--An iron
bottle, filled with iron borings, was heated in such a way that
the upper part became red hot, the lower part remaining cooler;
sodium hydroxide was introduced into the upper part. The
decomposition was then effected--that is, sodium vapours were
produced (this experiment was really performed with potassium
hydroxide). On opening the bottle it was found that the iron in
the upper part was not oxidised, but only that in the lower part.
This may be explained by the decomposition of the alkali into
sodium, hydrogen, and oxygen taking place in the upper part,
whilst the iron in the lower part absorbed the oxygen set free. If
the whole bottle be subjected to the same moderate heat as the
lower extremity, no metallic vapours are formed. In that case,
according to the hypothesis, the temperature is insufficient for
the dissociation of the sodium hydroxide.
[37] It has been previously remarked (Chapter II. Note 9) that Beketoff
showed the displacement of sodium by hydrogen, not from sodium
hydroxide but from the oxide Na_{2}O; then, however, only one half
is displaced, with the formation of NaHO.
For the production of sodium according to Deville's method, a mixture
of anhydrous sodium carbonate (7 parts), charcoal (two parts), and lime
or chalk (7 parts) is heated. This latter ingredient is only added in
order that the sodium carbonate, on fusing, shall not separate from
the charcoal.[38] The chalk on being heated loses carbonic anhydride,
leaving infusible lime, which is permeated by the sodium carbonate and
forms a thick mass, in which the charcoal is intimately mixed with the
sodium carbonate. When the charcoal is heated with the sodium carbonate,
at a white heat, carbonic oxide and vapours of sodium are disengaged,
according to the equation:
Na_{2}CO_{3} + 2C = Na_{2} + 3CO
[38] Since the close of the eighties in England, where the preparation
of sodium is at present carried out on a large commercial scale
(from 1860 to 1870 it was only manufactured in a few works in
France), it has been the practice to add to Deville's mixture
iron, or iron oxide which with the charcoal gives metallic and
carburetted iron, which still further facilitates the
decomposition. At present a kilogram of sodium may be purchased
for about the same sum (2/-) as a gram cost thirty years ago.
Castner, in England, greatly improved the manufacture of sodium in
large quantities, and so cheapened it as a reducing agent in the
preparation of metallic aluminium. He heated a mixture of 44 parts
of NaHO, and 7 parts of carbide of iron in large iron retorts at
1,000° and obtained about 6-1/2 parts of metallic sodium. The
reaction proceeds more easily than with carbon or iron alone, and
the decomposition of the NaHO proceeds according to the equation:
3NaHO + C = Na_{2}CO_{3} + 3H + Na. Subsequently, in 1891,
aluminium was prepared by electrolysis (_see_ Chapter XVII.), and
metallic sodium found two new uses; (1) for the manufacture of
peroxide of sodium (see later on) which is used in bleaching
works, and (2) in the manufacture of potassium and sodium cyanide
from yellow prussiate (Chapter XIII., Note 12).
[Illustration: FIG. 70.--Manufacture of sodium by Deville's process.
A C, iron tube containing a mixture of soda, charcoal, and chalk. B,
condenser.]
On cooling the vapours and gases disengaged, the vapours condense into
molten metal (in this form sodium does not easily oxidise, whilst in
vapour it burns) and the carbonic oxide remains as gas.
[Illustration: FIG. 71.--Donny and Maresca's sodium condenser, consisting
of two cast-iron plates screwed together.]
In sodium works an iron tube, about a metre long and a decimeter in
diameter, is made out of boiler plate. The pipe is luted into a furnace
having a strong draught, capable of giving a high temperature, and the
tube is charged with the mixture required for the preparation of sodium.
One end of the tube is closed with a cast-iron stopper A with clay
luting, and the other with the cast-iron stopper C provided with an
aperture. On heating, first of all the moisture contained in the various
substances is given off, then carbonic anhydride and the products of the
dry distillation of the charcoal, then the latter begins to act on the
sodium carbonate, and carbonic oxide and vapours of sodium appear. It is
easy to observe the appearance of the latter, because on issuing from
the aperture in the stopper C they take fire spontaneously and burn with
a very bright yellow flame. A pipe is then introduced into the aperture
C, compelling the vapours and gases formed to pass through the condenser
B. This condenser consists of two square cast-iron trays, A and A´, fig.
71, with wide edges firmly screwed together. Between these two trays
there is a space in which the condensation of the vapours of sodium is
effected, the thin metallic walls of the condenser being cooled by the
air but remaining hot enough to preserve the sodium in a liquid state,
so that it does not choke the apparatus, but continually flows from
it. The vapours of sodium, condensing in the cooler, flow in the shape
of liquid metal into a vessel containing some non-volatile naphtha or
hydrocarbon. This is used in order to prevent the sodium oxidising as
it issues from the condenser at a somewhat high temperature. In order
to obtain sodium of a pure quality it is necessary to distil it once
more, which may even be done in porcelain retorts, but the distillation
must be conducted in a stream of some gas on which sodium does not
act, for instance in a stream of nitrogen; carbonic anhydride is not
applicable, because sodium partially decomposes it, absorbing oxygen
from it. Although the above described methods of preparing sodium by
chemical means have proved very convenient in practice, still it is now
(since 1893) found profitable in England to obtain it (to the amount of
several tons a week) by Davy's classical method, _i.e._ by the action of
an electric current at a moderately high temperature, because the means
for producing an electric current (by motors and dynamos) now render
this quite feasible. This may be regarded as a sign that in process of
time many other technical methods for producing various substances by
_decomposition_ may be profitably carried on by electrolysis.
Pure sodium is a lustrous metal, white as silver, soft as wax; it
becomes brittle in the cold. In ordinary moist air it quickly tarnishes
and becomes covered with a film of hydroxide, NaHO, formed at the
expense of the water in the air. In perfectly dry air sodium retains its
lustre for an indefinite time. Its density at the ordinary temperature
is equal to 0·98, so that it is lighter than water; it fuses very
easily at a temperature of 95°, and distils at a bright red heat (742°
according to Perman, 1889). Scott (1887) determined the density of
sodium vapour and found it to be nearly 12 (if H = 1). This shows that
its molecule contains one atom (like mercury and cadmium) Na.[38 bis] It
forms alloys with most metals, combining with them, heat being sometimes
evolved and sometimes absorbed. Thus, if sodium (having a clean surface)
be thrown into mercury, especially when heated, there is a flash, and
such a considerable amount of heat is evolved that part of the mercury
is transformed into vapour.[39] Compounds or solutions of sodium in
mercury, or _amalgams_ of sodium, even when containing 2 parts of sodium
to 100 parts of mercury, are solids. Only those amalgams which are the
very poorest in sodium are liquid. Such alloys of sodium with mercury
are often used instead of sodium in chemical investigations, because in
combination with mercury sodium is not easily acted on by air, and is
heavier than water, and therefore more convenient to handle, whilst at
the same time it retains the principal properties of sodium,[40] for
instance it decomposes water, forming NaHO.
[38 bis] This is also shown by the fall in the temperature of
solidification of tin produced by the addition of sodium (and also
Al and Zn). Heycock and Neville (1889).
[39] By dissolving sodium amalgams in water and acids, and deducting
the heat of solution of the sodium, Berthelot found that _for each
atom of the sodium_ in amalgams containing a larger amount of
mercury than NaHg_{5}, the amount of heat evolved increases, after
which the heat of formation falls, and the heat evolved decreases.
In the formation of NaHg_{5} about 18,500 calories are evolved;
when NaHg_{3} is formed, about 14,000; and for NaHg about 10,000
calories. Kraft regarded the definite crystalline amalgam as
having the composition of NaHg_{6}, but at the present time, in
accordance with Grimaldi's results, it is thought to be NaHg_{5}.
A similar amalgam is very easily obtained if a 3 p.c. amalgam be
left several days in a solution of sodium hydroxide until a
crystalline mass is formed, from which the mercury may be removed
by strongly pressing in chamois leather. This amalgam with a
solution of potassium hydroxide forms a potassium amalgam,
KHg_{10}. It may be mentioned here that the latent heat of fusion
(of atomic quantities) of Hg = 360 (Personne), Na = 730 (Joannis),
and K = 610 calories (Joannis).
[40] Alloys are so similar to solutions (exhibiting such complete
parallelism in properties) that they are included in the same
class of so-called indefinite compounds. But in alloys, as
substances passing from the liquid to the solid state, it is
easier to discover the formation of definite chemical compounds.
Besides the alloys of Na with Hg, those with tin (Bailey 1892
found Na_{2}Sn), lead (NaPb), bismuth (Na_{3}Bi), &c. (Joannis
1892 and others) have been investigated.
It is easy to form an alloy of mercury and sodium having a crystalline
structure, and a definite atomic composition, NaHg_{5}. The alloy
of sodium with hydrogen or _sodium hydride_, Na_{2}H, which has the
external appearance of a metal,[41] is a most instructive example of
the characteristics of alloys. At the ordinary temperature sodium does
not absorb hydrogen, but from 300° to 421° the absorption takes place
at the ordinary pressure (and at an increased pressure even at higher
temperatures), as shown by Troost and Hautefeuille (1874). One volume of
sodium absorbs as much as 238 volumes of hydrogen. The metal increases
in volume, and when once formed the alloy can be preserved for some time
without change at the ordinary temperature. The appearance of sodium
hydride resembles that of sodium itself; it is as soft as this latter,
when heated it becomes brittle, and decomposes above 300°, evolving
hydrogen. In this decomposition all the phenomena of dissociation are
very clearly shown--that is, the hydrogen gas evolved has a definite
tension[42] corresponding with each definite temperature. This confirms
the fact that the formation of substances capable of dissociation can
only be accomplished within the dissociation limits. Sodium hydride
melts more easily than sodium itself, and then does not undergo
decomposition if it is in an atmosphere of hydrogen. It oxidises easily
in air, but not so easily as potassium hydride. The chemical reactions
of sodium are retained in its hydride, and, if we may so express it,
they are even increased by the addition of hydrogen. At all events,
in the properties of sodium hydride[43] we see other properties than
in such hydrogen compounds as HCl, H_{2}O, H_{3}N, H_{4}C, or even in
the gaseous metallic hydrides AsH_{3}, TeH_{2}. Platinum, palladium,
nickel, and iron, in absorbing hydrogen form compounds in which hydrogen
is in a similar state. In them, as in sodium hydride, the hydrogen is
compressed, absorbed, occluded (Chapter II.)[43 bis]
[41] Potassium forms a similar compound, but lithium, under the same
circumstances, does not.
[42] The tension of dissociation of hydrogen _p_, in millimetres of
mercury, is:--
_t_ = 330° 350° 400° 430°
for Na_{2}H _p_ = 28 57 447 910
for K_{2}H 45 72 548 1100
[43] In general, during the formation of alloys the volumes change
very slightly, and therefore from the volume of Na_{2}H some idea
may be formed of the volume of hydrogen in a solid or liquid
state. Even Archimedes concluded that there was gold in an alloy
of copper and gold by reason of its volume and density. From the
fact that the density of Na_{2}H is equal to 0·959, it may be seen
that the volume of 47 grams (the gram molecule) of this compound =
49·0 c.c. The volume of 46 grams of sodium contained in the
Na_{2}H (the density under the same conditions being 0·97) is
equal to 47·4 c.c. Therefore the volume of 1 gram of hydrogen in
Na_{2}H is equal to 1·6 c.c., and consequently the density of
metallic hydrogen, or the weight of 1 c.c., approaches 0·6 gram.
This density is also proper to the hydrogen alloyed with potassium
and palladium. Judging from the scanty information which is at
present available, liquid hydrogen near its absolute boiling point
(Chapter II.) has a much lower density.
[43 bis] We may remark that at low temperatures Na absorbs NH_{3} and
forms (NH_{3}Na)_{2} (_see_ Chapter VI., Note 14); this substance
absorbs CO and gives (NaCO)n (Chapter IX., Note 31), although by
itself Na does not combine directly with CO (but K does).
The most important chemical property of sodium is its power of easily
decomposing water and _evolving hydrogen_ from the majority of the
hydrogen compounds, and especially from all acids, and hydrates in which
hydroxyl must be recognised. This depends on its power of combining
with the elements which are in combination with the hydrogen. We
already know that sodium disengages hydrogen, not only from water,
hydrochloric acid,[44] and all other acids, but also from ammonia,[44 bis]
with the formation of sodamide NH_{2}Na, although it does not
displace hydrogen from the hydrocarbons.[45] Sodium burns both in
chlorine and in oxygen, evolving much heat. These properties are closely
connected with its power of taking up oxygen, chlorine, and similar
elements from most of their compounds. Just as it removes the oxygen
from the oxides of nitrogen and from carbonic anhydride, so also does
it decompose the majority of oxides at definite temperatures. Here the
action is essentially the same as in the decomposition of water. Thus,
for instance, when acting on magnesium chloride the sodium displaces the
magnesium, and when acting on aluminium chloride it displaces metallic
aluminium. Sulphur, phosphorus, arsenic and a whole series of other
elements, also combine with sodium.[46]
[44] H. A. Schmidt remarked that perfectly dry hydrogen chloride is
decomposed with great difficulty by sodium, although the
decomposition proceeds easily with potassium and with sodium in
moist hydrogen chloride. Wanklyn also remarked that sodium burns
with great difficulty in dry chlorine. Probably these facts are
related to other phenomena observed by Dixon, who found that
perfectly dry carbonic oxide does not explode with oxygen on
passing an electric spark.
[44 bis] Sodamide, NH_{2}Na, (Chapter IV., Note 14), discovered by
Gay-Lussac and Thénard, has formed the object of repeated
research, but has been most fully investigated by A. W. Titherley
(1894). Until recently the following was all that was known about
this compound:--
By heating sodium in dry ammonia, Gay-Lussac and Thénard obtained
an olive-green, easily-fusible mass, _sodamide_, NH_{2}Na,
hydrogen being separated. This substance with water forms sodium
hydroxide and ammonia; with carbonic oxide, CO, it forms sodium
cyanide, NaCN, and water, H_{2}O; and with dry hydrogen chloride
it forms sodium and ammonium chlorides. These and other reactions
of sodamide show that the metal in it preserves its energetic
properties in reaction, and that this compound of sodium is more
stable than the corresponding chlorine amide. When heated,
sodamide, NH_{2}Na, only partially decomposes, with evolution of
hydrogen, the principal part of it giving ammonia and sodium
nitride, Na_{3}N, according to the equation 3NH_{2}Na = 2NH_{3} +
NNa_{3}. The latter is an almost black powdery mass, decomposed by
water into ammonia and sodium hydroxide.
Titherley's researches added the following data:--
Iron or silver vessels should be used in preparing this body,
because glass and porcelain are corroded at 300°-400°, at which
temperature ammonia gas acts upon sodium and forms the amide with
the evolution of hydrogen. The reaction proceeds slowly, but is
complete if there be an excess of NH_{3}. Pure NH_{2}Na is
colourless (its colouration is due to various impurities),
semi-transparent, shows traces of crystallisation, has a
conchoidal fracture, and melts at 145°. Judging from the increase
in weight of the sodium and the quantity of hydrogen which is
disengaged, the composition of the amide is exactly NH_{2}Na. It
partially volatilises (sublimes) in vacuo at 200°, and breaks up
into 2Na + N_{2} + 2H_{2} at 500°. The same amide is formed when
oxide of sodium is heated in NH_{3}: Na_{2}O + 2NH_{3} = 2NaH_{2}N
+ H_{2}O. NaHO is also formed to some extent by the resultant
H_{2}O. Potassium and lithium form similar amides. With water,
alcohol, and acids, NH_{2}Na gives NH_{3} and NaHO, which react
further. Anhydrous CaO absorbs NH_{2}Na when heated without
decomposing it. When sodamide is heated with SiO_{2}, NH_{3} is
disengaged, and silicon nitride formed. It acts still more readily
upon boric anhydride when heated with it: 2NH_{2}Na + B_{2}O_{3} =
2BN + 2NaHO + H_{2}O. When slightly heated, NH_{2}Na + NOCl = NaCl
+ N_{2} + H_{2}O (NHNa_{2} and NNa_{3} are apparently not formed
at a higher temperature). The halogen organic compounds react with
the aid of heat, but with so much energy that the reaction
frequently leads to the ultimate destruction of the organic groups
and production of carbon.
[45] As sodium does not displace hydrogen from the hydrocarbons, _it
may be preserved_ in liquid hydrocarbons. Naphtha is generally
used for this purpose, as it consists of a mixture of various
liquid hydrocarbons. However, in naphtha sodium usually becomes
coated with a crust composed of matter produced by the action of
the sodium on certain of the substances contained in the mixture
composing naphtha. In order that sodium may retain its lustre in
naphtha, secondary octyl alcohol is added. (This alcohol is
obtained by distilling castor oil with caustic potash.) Sodium
keeps well in a mixture of pure benzene and paraffin.
[46] If sodium does not directly displace the hydrogen in hydrocarbons,
still by indirect means compounds may be obtained which contain
sodium and hydrocarbon groups. Some of these compounds have been
produced, although not in a pure state. Thus, for instance, zinc
ethyl, Zn(C_{2}H_{5})_{2}, when treated with sodium, loses zinc
and forms sodium ethyl, C_{2}H_{5}Na, but this decomposition is
not complete, and the compound formed cannot be separated by
distillation from the remaining zinc ethyl. In this compound the
energy of the sodium is clearly manifest, for it reacts with
substances containing haloids, oxygen, &c., and directly absorbs
carbonic anhydride, forming a salt of a carboxylic acid
(propionic).
With _oxygen_ sodium unites in three degrees of combination, forming
a suboxide Na_{4}O,[46 bis] an oxide, Na_{2}O, and a peroxide, NaO.
They are thus termed because Na_{2}O is a stable basic oxide (with
water it forms a basic hydroxide), whilst Na_{4}O and NaO do not
form corresponding saline hydrates and salts. The suboxide is a grey
inflammable substance which easily decomposes water, disengaging
hydrogen; it is formed by the slow oxidation of sodium at the ordinary
temperature. The peroxide is a greenish yellow substance, fusing at
a bright red heat; it is produced by burning sodium in an excess of
oxygen, and it yields oxygen when treated with water:
Suboxide: Na_{4}O + 3H_{2}O = 4NaHO + H_{2}[47]
Oxide: Na_{2}O + H_{2}O = 2NaHO[48]
Peroxide: Na_{2}O_{2} + H_{2}O = 2NaHO + O[49]
All three oxides form sodium hydroxide with water, but only the oxide
Na_{2}O is directly transformed into a hydrate. The other oxides liberate
either hydrogen or oxygen; they also present a similar distinction with
reference to many other agents. Thus carbonic anhydride combines directly
with the oxide Na_{2}O, which when heated in the gas burns, forming
sodium carbonate, whilst the peroxide yields oxygen in addition. When
treated with acids, sodium and all its oxides only form the salts
corresponding with sodium oxide--that is, of the formula or type NaX.
Thus the oxide of sodium, Na_{2}O, is _the only salt-forming oxide_ of
this metal, as water is in the case of hydrogen. Although the peroxide
H_{2}O_{2} is derived from hydrogen, and Na_{2}O_{2} from sodium, yet
there are no corresponding salts known, and if they are formed they are
probably as unstable as hydrogen peroxide. Although carbon forms carbonic
oxide, CO, still it has only one salt-forming oxide--carbonic anhydride,
CO_{2}. Nitrogen and chlorine both give several salt-forming oxides and
types of salts. But of the oxides of nitrogen, NO and NO_{2} do not form
salts, as do N_{2}O_{3}, N_{2}O_{4}, and N_{2}O_{5}, although N_{2}O_{4}
does not form special salts, and N_{2}O_{5} corresponds with the highest
form of the saline compounds of nitrogen. Such distinctions between the
elements, according to their power of giving one or several saline forms,
is a radical property of no less importance than the basic or acid
properties of their oxides. Sodium as a typical metal does not form any
acid oxides, whilst chlorine, as a typical non-metal, does not form bases
with oxygen. Therefore sodium _as an element_ may be thus characterised:
it forms one very stable salt-forming oxide, Na_{2}O, having powerful
basic properties, and its salts are of the general formula, NaX,
therefore in its compounds it is, like hydrogen, a basic and univalent
element.
[46 bis] It is even doubtful whether the suboxide exists
(_see_ Note 47).
[47] A compound, Na_{2}Cl, which corresponds with the suboxide, is
apparently formed when a galvanic current is passed through fused
common salt; the sodium liberated dissolves in the common salt,
and does not separate from the compound either on cooling or on
treatment with mercury. It is therefore supposed to be Na_{2}Cl;
the more so as the mass obtained gives hydrogen when treated with
water: Na_{2}Cl + H_{2}O = H + NaHO + NaCl, that is, it acts like
suboxide of sodium. If Na_{2}Cl really exists as a salt, then the
corresponding base Na_{4}O, according to the rule with other bases
of the composition M_{4}O, ought to be called a quaternary oxide.
According to certain evidence, a suboxide is formed when thin
sheets or fine drops of sodium slowly oxidise in moist air.
[48] According to observations easily made, sodium when fused in air
oxidises but does not burn, the combustion only commencing with
the formation of vapour--that is, when considerably heated. Davy
and Karsten obtained the oxides of potassium, K_{2}O, and of
sodium, Na_{2}O, by heating the metals with their hydroxides,
whence NaHO + Na = Na_{2}O + H, but N. N. Beketoff failed to
obtain oxides by this means. He prepared them by directly igniting
the metals in dry air, and afterwards heating with the metal in
order to destroy any peroxide. The oxide produced, Na_{2}O, when
heated in an atmosphere of hydrogen, gave a mixture of sodium and
its hydroxide: Na_{2}O + H = NaHO + Na (_see_ Chapter II., Note
9). If both the observations mentioned are accurate, then the
reaction is reversible. Sodium oxide ought to be formed during the
decomposition of sodium carbonate by oxide of iron (_see_ Note
26), and during the decomposition of sodium nitrite. According to
Karsten, its specific gravity is 2·8, according to Beketoff 2·3.
The difficulty in obtaining it is owing to an excess of sodium
forming the suboxide, and an excess of oxygen the peroxide. The
grey colour peculiar to the suboxide and oxide perhaps shows that
they contain metallic sodium. In addition to this, in the presence
of water it may contain sodium hydride and NaHO.
[49] Of the oxides of sodium, that easiest to form is the peroxide, NaO
or Na_{2}O_{2}; this is obtained when sodium is burnt in an excess
of oxygen. If NaNO_{3} be melted, it gives Na_{2}O_{2} with
metallic Na. In a fused state the peroxide is reddish yellow, but
it becomes almost colourless when cold. When heated with iodine
vapour, it loses oxygen: Na_{2}O_{2} +I_{2} = Na_{2}OI_{2} + O.
The compound Na_{2}OI_{2} is akin to the compound Cu_{2}OCl_{2}
obtained by oxidising CuCl. This reaction is one of the few in
which iodine directly displaces oxygen. The substance Na_{2}OI_{2}
is soluble in water, and when acidified gives free iodine and a
sodium salt. Carbonic oxide is absorbed by heated sodium peroxide
with formation of sodium carbonate: Na_{2}CO_{3} = Na_{2}O_{2} +
CO, whilst carbonic anhydride liberates oxygen from it. With
nitrous oxide it reacts thus: Na_{2}O_{2} +2N_{2}O = 2NaNO_{2}
+N_{2}; with nitric oxide it combines directly, forming sodium
nitrite, NaO + NO = NaNO_{2}. Sodium peroxide, when treated with
water, does not give hydrogen peroxide, because the latter in the
presence of the alkali formed (Na_{2}O_{2}+ 2H_{2}O = 2NaHO +
H_{2}O_{2}) decomposes into water and oxygen. In the presence of
dilute sulphuric acid it forms H_{2}O_{2} (Na_{2}O_{2} +
H_{2}SO_{4} = Na_{2}SO_{4} + H_{2}O_{2}). Peroxide of sodium is
now prepared on a large scale (by the action of air upon Na at
300°) for bleaching wool, silk &c. (when it acts in virtue of the
H_{2}O_{2} formed). The oxidising properties of Na_{2}O_{2} under
the action of heat are seen, for instance, in the fact that when
heated with I it forms sodium iodate; with PbO, Na_{2}PbO_{3};
with pyrites, sulphates, &c. When peroxide of sodium comes into
contact with water, it evolves much heat, forming H_{2}O_{2}, and
decomposing with the disengagement of oxygen; but, as a rule,
there is no explosion. But if Na_{2}O_{2} be placed in contact
with organic matter, such as sawdust, cotton, &c., it gives a
violent explosion when heated, ignited, or acted on by water.
Peroxide of sodium forms an excellent oxidising agent for the
preparation of the higher product of oxidation of Mn, Cr, W, &c.,
and also for oxidising the metallic sulphides. It should therefore
find many applications in chemical analysis. To prepare
Na_{2}O_{2} on a large scale, Castner melts Na in an aluminium
vessel, and at 300° passes first air deprived of a portion of its
oxygen (having been already once used), and then ordinary dry air
over it.
On comparing sodium and its analogues, which will be described later
with other metallic elements, it will be seen that these properties,
together with the relative lightness of the metal itself and its
compounds, and the magnitude of its atomic weight comprise the most
essential properties of this element, clearly distinguishing it from
others, and enabling us easily to recognise its analogues.
CHAPTER XIII
POTASSIUM, RUBIDIUM, CÆSIUM, AND LITHIUM. SPECTRUM ANALYSIS
Just as the series of halogens, fluorine, bromine and iodine correspond
with the chlorine contained in common salt, so also there exists a
corresponding series of elements: lithium, Li = 7, potassium, K = 39,
rubidium, Rb = 85, and cæsium, Cs = 133, which are analogous to the
sodium in common salt. These elements bear as great a resemblance to
sodium, Na = 23, as fluorine, F = 19, bromine, Br = 80, and iodine, I =
127, do to chlorine, Cl = 35·5. Indeed, in a free state, these elements,
like sodium, are soft metals which rapidly oxidise in moist air and
decompose water at the ordinary temperature, forming soluble hydroxides
having clearly-defined basic properties and the composition RHO, like
that of caustic soda. The resemblance between these metals is sometimes
seen with striking clearness, especially in compounds such as salts.[1]
The corresponding salts of nitric, sulphuric, carbonic, and nearly all
acids with these metals have many points in common. The metals which
resemble sodium so much in their reactions are termed the _metals of the
alkalis_.
[1] Tutton's researches (1894) upon the analogy of the crystalline
forms of K_{2}SO_{4}, Rb_{2}SO_{4} and Cs_{2}SO_{4} may be taken
as a typical example of the comparison of analogous compounds. We
cite the following data from these excellent researches: the sp.
gr. at 20°/4° of K_{2}SO_{4} is 2·6633 of Rb_{2}SO_{4}, 3·6113, and
of Cs_{2}SO_{4}, 4·2434. The coefficient of cubical expansion (the
mean between 20° and 60°) for the K salt is 0·0053, for the Rb salt
0·0052, for the Cs salt 0·0051. The linear expansion (the maximum
for the vertical axis) along the axis of crystallisation is the
same for all three salts, within the limits of experimental error.
The replacement of potassium by rubidium causes the distance
between the centres of the molecules in the direction of the three
axes of crystallisation to increase equally, and less than with the
replacement of rubidium by cæsium. The index of refraction for all
rays and for every crystalline path (direction) is greater for the
rubidium salt than for the potassium salt, and less than for the
cæsium salt, and the differences are nearly in the ratio 2 : 5. The
lengths of the rhombic crystalline axes for K_{2}SO_{4} are in the
ratio 0·5727 : 1 : 0·7418, for Rb_{2}SO_{4}, 0·5723 : 1 : 0·7485,
and for Cs_{2}SO_{4}, 0·5712 : 1 : 0·7521. The development of the
basic and brachy-pinacoids gradually increases in passing from K to
Rb and Cs. The optical properties also follow the same order both
at the ordinary and at a higher temperature. Tutton draws the
general conclusion that the crystallographic properties of the
isomorphic rhombic sulphates R_{2}SO_{4} are a function of the
atomic weight of the metals contained in them (_see_ Chapter XV.)
Such researches as these should do much towards hastening the
establishment of a true molecular mechanics of physico-chemical
phenomena.
Among the metals of the alkalis, the most widely distributed in nature,
after sodium, is _potassium_. Like sodium, it does not appear either in
a free state or as oxide or hydroxide, but in the form of salts, which
present much in common with the salts of sodium in the manner of their
occurrence. The compounds of potassium and sodium in the earth's crust
occur as mineral compounds of silica. With silica, SiO_{2}, potassium
oxide, like sodium oxide, forms saline mineral substances resembling
glass. If other oxides, such as lime, CaO, and alumina, Al_{2}O_{3},
combine with these compounds, glass is formed, a vitreous stony mass,
distinguished by its great stability, and its very slight variation
under the action of water. It is such complex silicious compounds as
these which contain potash (potassium oxide), K_{2}O, or soda (sodium
oxide), Na_{2}O, and sometimes both together, silica, SiO_{2}, lime,
CaO, alumina, Al_{2}O_{3}, and other oxides, that form the chief mass
of rocks, out of which, judging by the direction of the strata, the
chief mass of the accessible crust (envelope) of the earth is made up.
The primary rocks, like granite, porphyry, &c.,[1 bis] are formed of
such crystalline silicious rocks as these. The oxides entering into the
composition of these rocks do not form a homogeneous amorphous mass
like glass, but are distributed in a series of peculiar, and in the
majority of cases crystalline, compounds, into which the primary rocks
may be divided. Thus a felspar (orthoclase) in granite contains from
8 to 15 per cent. of potassium, whilst another variety (plagioclase)
which also occurs in granite contains 1·2 to 6 per cent. of potassium,
and 6 to 12 per cent. of sodium. The mica in granite contains 3 to 10
per cent. of potassium. As already mentioned, and further explained in
Chapter XVII., the friable, crumbling, and stratified formations which
in our times cover a large part of the earth's surface have been formed
from these primary rocks by the action of the atmosphere and of water
containing carbonic acid. It is evident that in the chemical alteration
of the primary rocks by the action of water, the compounds of potassium,
as well as the compounds of sodium, must have been dissolved by the
water (as they are soluble in water), and that therefore the compounds
of potassium must be accumulated together with those of sodium in sea
water. And indeed compounds of potassium are always found in _sea
water_, as we have already pointed out (Chapters I. and X.). This forms
one of the sources from which they are extracted. After the evaporation
of sea water, there remains a mother liquor, which contains potassium
chloride and a large proportion of magnesium chloride. On cooling this
solution crystals separate out which contain chlorides of magnesium
and potassium. A double salt of this kind, called _carnallite_,
KMgCl_{3},6H_{2}O, occurs at Stassfurt. This carnallite[2] is now
employed as a material for the extraction of potassium chloride, and
of all the compounds of this element.[3] Besides which, potassium
chloride itself is sometimes found at Stassfurt as _sylvine_.[3 bis]
By a method of double saline decomposition, the chloride of potassium
may be converted into all the other potassium salts,[4] some of which
are of practical use. The potassium salts have, however, their greatest
importance as an indispensable component of the food of plants.[5]
[1 bis] The origin of the primary rocks has been mentioned in
Chapter X., Note 2.
[2] Carnallite belongs to the number of double salts which are directly
decomposed by water, and it only crystallises from solutions which
contain an excess of magnesium chloride. It may be prepared
artificially by mixing strong solutions of potassium and magnesium
chlorides, when colourless crystals of sp. gr. 1·60 separate,
whilst the Stassfurt salt is usually of a reddish tint, owing to
traces of iron. At the ordinary temperature sixty-five parts of
carnallite are soluble in one hundred parts of water in the
presence of an excess of the salt. It deliquesces in the air,
forming a solution of magnesium chloride and leaving potassium
chloride. The quantity of carnallite produced at Stassfurt is now
as much as 100,000 tons a year.
[3] The method of separating sodium chloride from potassium chloride
has been described in Chapter I. On evaporation of a mixture of the
saturated solutions, sodium chloride separates; and then, on
cooling, potassium chloride separates, owing to the difference of
rate of variation of their solubilities with the temperature. The
following are the most trustworthy figures for the solubility of
_potassium chloride_ in one hundred parts of water (for sodium
chloride, _see_ Chapter X., Note 13):--
10° 20° 40° 60° 100°
32 35 40 46 57
When mixed with solutions of other salts the solubility of
potassium chloride naturally varies, but not to any great extent.
[3 bis] The specific gravity of the solid salt is 1·99--that is, less
than that of sodium chloride. All the salts of sodium are
specifically heavier than the corresponding salts of potassium, as
are also their solutions for equal percentage compositions. If the
specific gravity of water at 4° = 10,000, then at 15° the specific
gravity of a solution of _p_ p.c. potassium chloride = 9,992 +
63·29_p_ + 0·226_p_^2, and therefore for 10 p.c. = 1·0647, 20 p.c.
= 1·1348, &c.
Potassium chloride combines with iodine trichloride to form a
compound KCl + ICl_{3} = KICl_{4}, which has a yellow colour, is
fusible, loses iodine trichloride at a red heat, and gives
potassium iodate and hydrochloric acid with water. It is not only
formed by direct combination, but also by many other methods; for
instance, by passing chlorine into a solution of potassium iodide
so long as the gas is absorbed, KI + 2Cl_{2} = KCl,ICl_{3}.
Potassium iodide, when treated with potassium chlorate and strong
hydrochloric acid, also gives this compound; another method for its
formation is given by the equation KClO_{3} + I + 6HCl =
KCl,ICl_{3} + 3Cl + 3H_{2}O. This is a kind of salt corresponding
with KIO_{2} (unknown) in which the oxygen is replaced by chlorine.
If valency be taken as the starting-point in the study of chemical
compounds, and the elements considered as having a constant
atomicity (number of bonds)--that is, if K, Cl, and I be taken as
univalent elements--then it is impossible to explain the formation
of such a compound because, according to this view, univalent
elements are only able to form dual compounds with each other; such
as, KCl, ClI, KI, &c., whilst here they are grouped together in the
molecule KICl_{4}. Wells, Wheeler, and Penfield (1892) obtained a
large number of such poly-haloid salts. They may all be divided
into two large classes: the tri-haloid and the penta-haloid salts.
They have been obtained not only for K but also for Rb and Cs, and
partially also for Na and Li. The general method of their formation
consists in dissolving the ordinary halogen salt of the metal in
water, and treating it with the requisite amount of free halogen.
The poly-haloid salt separates out after evaporating the solution
at a more or less low temperature. In this manner, among the
tri-haloid salts, may be obtained: KI_{3}, KBr_{2}I, KCl_{2}I, and
the corresponding salts of rubidium and cæsium, for instance,
CsI_{3}, CsBrI_{2}, CsBr_{2}I, CsClBrI,CsCl_{2}I, CsBr_{3},
CsClBr_{2}, CsCl_{2}Br, and in general MX_{3} where X is a halogen.
The colour of the crystals varies according to the halogen, thus
CsI_{3} is black, CrBr_{3} yellowish red, CrBrI_{2} reddish brown,
CsBr_{2}I red, CsCl_{2}Br yellow. The cæsium salts are the most
stable, and those of potassium least so, as also those which
contain Br and I separately or together; for cæsium no compounds
containing Cl and I were obtained. The penta-haloid salts form a
smaller class; among these salts potassium forms KCl_{4}I, rubidium
RbCl_{4}I, cæsium CsI_{5}, CsBr, CsCl_{4}I, lithium LiCl_{4}I (with
4H_{2}O) and sodium NaCl_{4}I (with 2H_{2}O). The most stable are
those salts containing the metal with the greatest atomic
weight--cæsium (_see_ Chapter XI., Note 63).
[4] It is possible to extract the compounds of potassium directly from
the primary rocks which are so widely distributed over the earth's
surface and so abundant in some localities. From a chemical point
of view this problem presents no difficulty; for instance, by
fusing powdered orthoclase with lime and fluor spar (Ward's method)
and then extracting the alkali with water (on fusion the silica
gives an insoluble compound with lime), or by treating the
orthoclase with hydrofluoric acid (in which case silicon fluoride
is evolved as a gas) it is possible to transfer the alkali of the
orthoclase to an aqueous solution, and to separate it in this
manner from the other insoluble oxides. However, as yet there is no
profit in, nor necessity for, recourse to this treatment, as
carnallite and potash form abundant materials for the extraction of
potassium compounds by cheaper methods. Furthermore, the salts of
potassium are now in the majority of chemical reactions replaced by
salts of sodium, especially since the preparation of sodium
carbonate has been facilitated by the Leblanc process. The
replacement of potassium compounds by sodium compounds not only has
the advantage that the salts of sodium are in general cheaper than
those of potassium, but also that a smaller quantity of a sodium
salt is needed for a given reaction than of a potassium salt,
because the combining weight of sodium (23) is less than that of
potassium (39).
[5] It has been shown by direct experiment on the cultivation of plants
in artificial soils and in solutions that under conditions
(physical, chemical, and physiological) otherwise identical plants
are able to thrive and become fully developed in the entire absence
of sodium salts, but that their development is impossible without
potassium salts.
The primary rocks contain an almost equal proportion of potassium and
sodium. But in sea water the compounds of the latter metal predominate.
It may be asked, what became of the compounds of potassium in the
disintegration of the primary rocks, if so small a quantity went to the
sea water? They remained with the other products of the decomposition
of the primary rocks. When granite or any other similar rock formation
is disintegrated, there are formed, besides the soluble substances,
also insoluble substances--sand and finely-divided clay, containing
water, alumina, and silica. This clay is carried away by the water,
and is then deposited in strata. It, and especially its admixture with
vegetable remains, retain compounds of potassium in a greater quantity
than those of sodium. This has been proved with absolute certainty
to be the case, and is due to the _absorptive power of the soil_. If
a dilute solution of a potassium compound be filtered through common
mould used for growing plants, containing clay and the remains of
vegetable decomposition, this mould will be found to have retained a
somewhat considerable percentage of the potassium compounds. If a salt
of potassium be taken, then during the filtration an equivalent quantity
of a salt of calcium--which is also found, as a rule, in soils--is
set free. Such a process of filtration through finely divided earthy
substances proceeds in nature, and the compounds of potassium are
everywhere retained by the friable earth in considerable quantity. This
explains the presence of so small an amount of potassium salts in the
water of rivers, lakes, streams, and oceans, where the lime and soda
have accumulated. The compounds of potassium retained by the friable
mass of the earth are absorbed as an aqueous solution by the roots
of _plants_. Plants, as everyone knows, when burnt leave an ash, and
this ash, besides various other substances, without exception contains
compounds of potassium. Many land plants contain a very small amount of
sodium compounds,[6] whilst potassium and its compounds occur in all
kinds of vegetable ash. Among the generally cultivated plants, grass,
potatoes, the turnip, and buckwheat are particularly rich in potassium
compounds. The ash of plants, and especially of herbaceous plants,
buckwheat straw, sunflower and potato leaves are used in practice for
the extraction of potassium compounds. There is no doubt that potassium
occurs in the plants themselves in the form of complex compounds,
and often as salts of organic acids. In certain cases such salts of
potassium are even extracted from the juice of plants. Thus, sorrel
and oxalis, for example, contain in their juices the acid oxalate of
potassium, C_{2}HKO_{4}, which is employed for removing ink stains.
Grape juice contains the so-called cream of tartar, which is the acid
tartrate of potassium, C_{4}H_{5}KO_{6}.[7] This salt also separates
as a sediment from wine. When the plants, containing one or more of the
salts of potassium, are burnt, the carbonaceous matter is oxidised,
and in consequence the potassium is obtained in the ash as carbonate,
K_{2}CO_{3}, which is generally known as _potashes_. Hence potashes
occur ready prepared in the ash of plants, and therefore the ash of
land plants is employed as a source for the extraction of potassium
compounds. Potassium carbonate is extracted by lixiviating the ash with
water.[8] Potassium carbonate may also be obtained from the chloride
by a method similar to that by which sodium carbonate is prepared from
sodium chloride.[8 bis] There is no difficulty in obtaining any salt of
potassium--for example, the sulphate,[9] bromide, and iodide[10]--by
the action of the corresponding acid on KCl and especially on the
carbonate, whilst the hydroxide, _caustic potash_, KHO, which is in many
respects analogous to caustic soda, is easily obtained by means of lime
in exactly the same manner in which sodium hydroxide is prepared from
sodium carbonate.[11] Therefore, in order to complete our knowledge of
the alkali metals, we will only describe two salts of potassium which
are of practical importance, and whose analogues have not been described
in the preceding chapter, potassium cyanide and potassium nitrate.
[6] If herbaceous plants contain much sodium salts, it is evident that
these salts mainly come from the sodium compounds in the water
absorbed by the plants.
[7] As plants always contain mineral substances and cannot thrive in
a medium which does not contain them, more especially in one which
is free from the salts of the four basic oxides, K_{2}O, CaO, MgO,
and Fe_{2}O_{3}, and of the four acid oxides, CO_{2}, N_{2}O_{5},
P_{2}O_{5}, and SO_{3}, and as the amount of ash-forming substances
in plants is small, the question inevitably arises as to what part
these play in the development of plants. With the existing chemical
data only one answer is possible to this question, and it is still
only a hypothesis. This answer was particularly clearly expressed
by Professor Gustavson of the Petroffsky Agricultural Academy.
Starting from the fact (Chapter XI., Note 55) that a small quantity
of aluminium renders possible or facilitates the reaction of
bromine on hydrocarbons at the ordinary temperature, it is easy to
arrive at the conclusion, which is very probable and in accordance
with many data respecting the reactions of organic compounds, that
the addition of mineral substances to organic compounds lowers the
temperature of reaction and in general facilitates chemical
reactions in plants, and thus aids the conversion of the most
simple nourishing substances into the complex component parts of
the plant organism. The province of chemical reactions proceeding
in organic substances in the presence of a small quantity of
mineral substances has as yet been but little investigated,
although there are already several disconnected data concerning
reactions of this kind, and although a great deal is known with
regard to such reactions among inorganic compounds. The essence of
the matter may be expressed thus--two substances, A and B, do not
react on each other of their own accord, but the addition of a
small quantity of a third particularly active substance, C,
produces the reaction of A on B, because A combines with C, forming
AC, and B reacts on this new compound, which has a different store
of chemical energy, forming the compound AB or its products, and
setting C free again or retaining it.
It may here be remarked that all the mineral substances necessary
for plants (those enumerated at the beginning of the note) are the
highest saline compounds of their elements, that they enter into
the plants as salts, that the lower forms of oxidation of the same
elements (for instance, sulphites and phosphites) are harmful to
plants (poisonous), and that strong solutions of the salts
assimilated by plants (their osmotic pressure being great and
contracting the cells, as De Vries showed, (_see_ Chapter I., Note
19)) not only do not enter into the plants but kill them (poison
them).
[8] Besides which, it will be understood from the preceding paragraph
that the salts of potassium may become exhausted from the soil by
long cultivation, and that there may therefore be cases when the
direct fertilisation by salts of potassium may be profitable. But
manure and animal excrements, ashes, and, in general, nearly all
refuse which may serve for fertilising the soil, contain a
considerable quantity of potassium salts, and therefore, as regards
the natural salts of potassium (Stassfurt), and especially
potassium sulphate, if they often improve the crops, it is in all
probability due to their action on the properties of the soil. The
agriculturist cannot therefore be advised to add potassium salts,
without making special experiments showing the advantage of such a
fertiliser on a given kind of soil and plant.
The animal body also contains potassium compounds, which is
natural, since animals consume plants. For example, milk, and
especially human milk, contains a somewhat considerable quantity of
potassium compounds. Cow's milk, however, does not contain much
potassium salt. Sodium compounds generally predominate in the
bodies of animals. The excrement of animals, and especially of
herbivorous animals, on the contrary, often contains a large
proportion of potassium salts. Thus sheep's dung is rich in them,
and in washing sheep's wool salts of potassium pass into the water.