Chapter I: , because in many instances unstable highly iodised (2)

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.

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]

_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:--