Chapter I: , because in many instances unstable highly iodised (3)

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).

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.

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.