Thomson concludes his investigation with the words: (_a_) 'When

equivalent quantities of NaHO, HNO_{3} (or HCl) and 1/2H_{2}SO_{4}

react on one another in an aqueous solution, then two-thirds of

the soda combines with the nitric and one-third with the sulphuric

acid; (_b_) this subdivision repeats itself, whether the soda be

taken combined with nitric or with sulphuric acid; (_c_) and

therefore nitric acid has double the tendency to combine with the

base that sulphuric acid has, and hence in an aqueous solution it

is a stronger acid than the latter.'

'It is therefore necessary,' Thomsen afterwards remarks, 'to have

an expression indicating the tendency of an acid for the

saturation of bases. This idea cannot be expressed by the word

_affinity_, because by this term is most often understood that

force which it is necessary to overcome in order to decompose a

substance into its component parts. This force should therefore be

measured by the amount of work or heat employed for the

decomposition of the substance. The above-mentioned phenomenon is

of an entirely different nature,' and Thomsen introduces the term

_avidity_, by which he designates the tendency of acids for

neutralisation. 'Therefore the avidity of nitric acid with respect

to soda is twice as great as the avidity of sulphuric acid. An

exactly similar result is obtained with hydrochloric acid, so that

its avidity with respect to soda is also double the avidity of

sulphuric acid. Experiments conducted with other acids showed that

not one of the acids investigated had so great an avidity as

nitric acid; some had a greater avidity than sulphuric acid,

others less, and in some instances the avidity = 0.' The reader

will naturally see clearly that the path chosen by Thomsen

deserves to be worked out, for his results concern important

questions of chemistry, but great faith cannot be placed in the

deductions he has already arrived at, because great complexity of

relations is to be seen in the very method of his investigation.

It is especially important to turn attention to the fact that all

the reactions investigated are reactions of double decomposition.

In them A and B do not combine with C and distribute themselves

according to their affinity or avidity for combination, but

reversible reactions are induced. MX and NY give MY and NX, and

conversely; therefore the affinity or avidity for combination is

not here directly determined, but only the difference or relation

of the affinities or avidities. The affinity of nitric acid not

only for the water of constitution, but also for that serving for

solution, is much less than that of sulphuric acid. This is seen

from thermal data. The reaction N_{2}O_{5} + H_{2}O gives +3,600

heat units, and the solution of the resultant hydrate, 2NHO_{3},

in a large excess of water evolves +14,986 heat units. The

formation of SO_{3} + H_{2}O evolves +21,308 heat units, and the

solution of H_{2}SO_{4} in an excess of water 17,860--that is,

sulphuric acid gives more heat in both cases. The interchange

between Na_{2}SO_{4} and 2HNO_{3} is not only accomplished at the

expense of the production of NaNO_{3}, but also at the expense of

the formation of H_{2}SO_{4}, hence the affinity of sulphuric acid

for water plays its part in the phenomena of displacement.

Therefore in determinations like those made by Thomsen the water

does not form a medium which is present without participating in

the process; it also takes part in the reaction. (Compare Chapter

IX., Note 14.)

Whilst retaining essentially the methods of Thomsen, Ostwald

(1876) determined the variation of the sp. gr. (and afterwards of

volume), proceeding in the same dilute solutions, on the

saturation of acids by bases, and in the decomposition of the

salts of one acid by the other, and arrived at conclusions of just

the same nature as Thomsen's. Ostwald's method will be clearly

understood from an example. A solution of caustic soda containing

an almost molecular (40 grams) weight per litre had a specific

gravity of 1·04051. The specific gravities of solutions of equal

volume and equivalent composition of sulphuric and nitric acids

were 1·02970 and 1·03084 respectively. On mixing the solutions of

NaHO and H_{2}SO_{4} there was formed a solution of Na_{2}SO_{4}

of sp. gr. 1·02959; hence there ensued a decrease of specific

gravity which we will term Q, equal to 1·04051 +

1·02970-2(1·02959) = 0·01103. So also the specific gravity after

mixture of the solutions of NaHO and HNO_{3} was 1·02633, and

therefore Q = 0·01869. When one volume of the solution of nitric

acid was added to two volumes of the solution of sodium sulphate,

a solution of sp. gr. 1·02781 was obtained, and therefore the

resultant decrease of sp. gr.

Q_{1} = 2(1·02959) + 1·03084-3(1·02781) = 0·00659.

Had there been no chemical reaction between the salts, then

according to Ostwald's reasoning the specific gravity of the

solutions would not have changed, and if the nitric acid had

entirely displaced the sulphuric acid Q_{2} would be =

0·01869-0·01103 = 0·00766. It is evident that a portion of the

sulphuric acid was displaced by the nitric acid. But the measure

of displacement is not equal to the ratio between Q_{1} and Q_{2},

because a decrease of sp. gr. also occurs on mixing the solution

of sodium sulphate with sulphuric acid, whilst the mixing of the

solutions of sodium nitrate and nitric acid only produces a slight

variation of sp. gr. which falls within the limits of experimental

error. Ostwald deduces from similar data the same conclusions as

Thomsen, and thus reconfirms the formula deduced by Guldberg and

Waage, and the teaching of Berthollet.

The participation of water is seen still more clearly in the

methods adopted by Ostwald than in those of Thomsen, because in

the saturation of solutions of acids by alkalis (which Kremers,

Reinhold, and others had previously studied) there is observed,

not a contraction, as might have been expected from the quantity

of heat which is then evolved, but an expansion, of volume (a

decrease of specific gravity, if we calculate as Ostwald did in

his first investigations). Thus by mixing 1,880 grams of a

solution of sulphuric acid of the composition SO_{3} + 100H_{2}O,

occupying a volume of 1,815 c.c., with a corresponding quantity of

a solution 2(NaHO + 5H_{2}O), whose volume = 1,793 c.c., we obtain

not 3,608 but 3,633 c.c., an expansion of 25 c.c. per gram

molecule of the resulting salt, Na_{2}SO_{4}. It is the same in

other cases. Nitric and hydrochloric acids give a still greater

expansion than sulphuric acid, and potassium hydroxide than sodium

hydroxide, whilst a solution of ammonia gives a contraction. The

relation to water must be considered as the cause of these

phenomena. When sodium hydroxide and sulphuric acid dissolve in

water they develop heat and give a vigorous contraction; the water

is separated from such solutions with great difficulty. After

mutual saturation they form the salt Na_{2}SO_{4}, which retains

the water but feebly and evolves but little heat with it, i.e., in

other words, has little affinity for water. In the saturation of

sulphuric acid by soda the water is, so to say, displaced from a

stable combination and passes into an unstable combination; hence

an expansion (decrease of sp. gr.) takes place. It is not the

reaction of the acid on the alkali, but the reaction of water,

that produces the phenomenon by which Ostwald desires to measure

the degree of salt formation. The water, which escaped attention,

itself has affinity, and influences those phenomena which are

being investigated. Furthermore, in the given instance its

influence is very great because its mass is large. When it is not

present, or only present in small quantities, the attraction of

the base to the acid leads to contraction, and not expansion.

Na_{2}O has a sp. gr. 2·8, hence its molecular volume = 22; the

sp. gr. of SO_{3} is 1·9 and volume 41, hence the sum of their

volumes is 63; for Na_{2}SO_{4} the sp. gr. is 2·65 and volume

53·6, consequently there is a contraction of 10 c.c. per

gram-molecule of salt. The volume of H_{2}SO_{4} = 53·3, that of

2NaHO = 37·4; there is produced 2H_{2}O, volume = 36, +

Na_{2}SO_{4}, volume = 53·6. There react 90·7 c.c., and on

saturation there result 89·6 c.c.; consequently contraction again

ensues, although less, and although this reaction is one of

substitution and not of combination. Consequently the phenomena

studied by Ostwald depend but little on the measure of the

reaction of the salts, and more on the relations of the dissolved

substances to water. In substitutions, for instance 2NaNO_{3} +

H_{2}SO_{4} = 2HNO_{3} + Na_{2}SO_{4}, the volumes vary but

slightly: in the above example they are 2(38·8) + 53·3 and 2(41·2)

+ 53·6; hence 131 volumes act, and 136 volumes are produced. It

may be concluded, therefore, on the basis of what has been said,

that on taking water into consideration the phenomena studied by

Thomsen and Ostwald are much more complex than they at first

appear, and that this method can scarcely lead to a correct

interpretation as to the distribution of acids between bases. We

may add that P. D. Chroustcheff (1890) introduced a new method for

this class of research, by investigating the electro-conductivity

of solutions and their mixtures, and obtained remarkable results

(for example, that hydrochloric acid almost entirely displaces

formic acid and only 2/3 of sulphuric acid), but details of these

methods must be looked for in text-books of theoretical

chemistry.

[28] G. G. Gustavson's researches, which were conducted in the

laboratory of the St. Petersburg University in 1871-72, are among

the first in which the measure of the affinity of the elements for

the halogens is recognised with perfect clearness in the limit of

substitution and in the rate of reaction. The researches conducted

by A. L. Potilitzin (of which mention will be made in Chapter XI.,

Note 66) in the same laboratory touch on another aspect of the

same problem which has not yet made much progress, notwithstanding

its importance and the fact that the theoretical side of the

subject (thanks especially to Guldberg and Van't Hoff) has since

been rapidly pushed forward. If the researches of Gustavson took

account of the influence of mass, and were more fully supplied

with data concerning velocities and temperatures, they would be

very important, because of the great significance which the case

considered has for the understanding of double saline

decompositions in the absence of water.

Furthermore Gustavson showed that the greater the atomic weight of

the element (B, Si, Ti, As, Sn) combined _with chlorine_ the

greater the amount of chlorine replaced by bromine by the action

of CBr_{4}, and consequently the less the amount of bromine

replaced by chlorine by the action of CCl_{4} on bromine

compounds. For instance, for chlorine compounds the percentage of

substitution (at the limit) is--

BCl_{3} SiCl_{4} TiCl_{4} AsCl_{3} SnCl_{4}

10·1 12·5 43·6 71·8 77·5

It should he observed, however, that Thorpe, on the basis of his

experiments, denies the universality of this conclusion. I may

mention one conclusion which it appears to me may be drawn from

the above-cited figures of Gustavson, if they are subsequently

verified even within narrow limits. If CBr_{4} be heated with

RCl_{4}, then an exchange of the bromine for chlorine takes place.

But what would be the result if it were mixed with CCl_{4}?

Judging by the magnitude of the atomic weights, B = 11, C = 12, Si

= 28, about 11 p.c. of the chlorine would be replaced by bromine.

But to what does this point? I think that this shows the existence

of a motion of the atoms in the molecule. The mixture of CCl_{4}

and CBr_{4} does not remain in a condition of static equilibrium;

not only are the molecules contained in it in a state of motion,

but also the atoms in the molecules, and the above figures show

the measure of their translation under these conditions. The

bromine in the CBr_{4} is, _within the limit_, substituted by the

chlorine of the CCl_{4} in a quantity of about 11 out of 100: that

is, a portion of the atoms of bromine previously to this moment in

combination with one atom of carbon pass over to the other atom of

carbon, and the chlorine passes over from this second atom of

carbon to replace it. Therefore, also, in the homogeneous mass

CCl_{4} all the atoms of Cl do not remain constantly combined with

the same atoms of carbon, and _there is on exchange of atoms

between different molecules in a homogeneous medium also_. This

hypothesis may in my opinion explain certain phenomena of

dissociation, but though mentioning it I do not consider it worth

while to dwell upon it. I will only observe that a similar

hypothesis suggested itself to me in my researches on solutions,

and that Pfaundler enunciated an essentially similar hypothesis,

and in recent times a like view is beginning to find favour with

respect to the electrolysis of saline solutions.

Thus we now find ample confirmation from various quarters for the

following rules of Berthollet, applying them to double saline

decompositions: 1. From two salts MX and NY containing different haloids

and metals there result from their reaction two others, MY and NX, but

such a substitution will not proceed to the end unless one product passes

from the sphere of action. 2. This reaction is limited by the existence

of an equilibrium between MX, NY, MY, and NX, because a reverse reaction

is quite as possible as the direct reaction. 3. This limit is determined

both by the measure of the active affinities and by the relative masses

of the substances as measured by the number of the reacting molecules. 4.

Other conditions being constant, the chemical action is proportional to

the product of the chemical masses in action.[29]

[29] Berthollet's doctrine is hardly at all affected in principle by

showing that there are cases in which there is no decomposition

between salts, because the affinity may be so small that even a

large mass would still give no observable displacements. The

fundamental condition for the application of Berthollet's

doctrine, as well as Deville's doctrine of dissociation, lies in

the reversibility of reactions. There are practically irreversible

reactions (for instance, CCl_{4} + 2H_{2}O = CO_{2} + 4HCl), just

as there are non-volatile substances. But while accepting the

doctrine of reversible reactions and retaining the theory of the

evaporation of liquids, it is possible to admit the existence of

non-volatile substances, and in just the same way of reactions,

without any visible conformity to Berthollet's doctrine. This

doctrine evidently comes nearer than the opposite doctrine of

Bergmann to solving the complex problems of chemical mechanics for

the successful solution of which at the present time the most

valuable help is to be expected from the working out of data

concerning dissociation, the influence of mass, and the

equilibrium and velocity of reactions. But it is evident that from

this point of view we must not regard a solvent as a

non-participant space, but must take into consideration the

chemical reactions accompanying solution, or else bring about

reactions without solution.

Thus if the salts MX and NY after reaction partly formed salts MY and NX,

then a state of equilibrium is reached and the reaction ceases; but if

one of the resultant compounds, in virtue of its physical properties,

passes from the sphere of action of the remaining substances, then the

reaction will continue. This exit from the sphere of action depends on

the physical properties of the substance and on the conditions under

which the reaction takes place. Thus, for instance, the salt NX may, in

the case of reaction between solutions, separate as a precipitate, an

insoluble substance, while the other three substances remain in solution,

or it may pass into vapour, and in this manner also pass away from the

sphere of action of the remaining substances. Let us now suppose that it

passes away in some form or other from the sphere of action of the

remaining substances--for instance, that it is transformed into a

precipitate or vapour--then a fresh reaction will set in and a

re-formation of the salt NX. If this be removed, then, although the

quantity of the elements N and X in the mass will be diminished, still,

according to Berthollet's law, a certain amount of NX should be again

formed. When this substance is again formed, then, owing to its physical

properties, it will again pass away; hence the reaction, in consequence

of the physical properties of the resultant substances, is able to

proceed to completion notwithstanding the possible weakness of the

attraction existing between the elements entering into the composition of

the resultant substance NX. Naturally, if the resultant substance is

formed of elements having a considerable degree of affinity, then the

complete decomposition is considerably facilitated.

Such a representation of the _modus operandi_ of chemical

transformations is applicable with great clearness to a number of

reactions studied in chemistry, and, what is especially important, the

application of this aspect of Berthollet's teaching does not in any way

require the determination of the measure of affinity acting between the

substances present. For instance, the action of ammonia on solutions of

salts; the displacement, by its means, of basic hydrates insoluble in

water; the separation of volatile nitric acid by the aid of non-volatile

sulphuric acid, as well as the decomposition of common salt by means of

sulphuric acid, when gaseous hydrochloric acid is formed--may be taken

as examples of reactions which proceed to the end, inasmuch as one of

the resultant substances is entirely removed from the sphere of action,

but they in no way indicate the measure of affinity.[30]

[30] Common salt not only enters into double decomposition with acids

but also _with every salt_. However, as clearly follows from

Berthollet's doctrine, this form of decomposition will only in a

few cases render it possible for new metallic chlorides to be

obtained, because the decomposition will not be carried on to the

end unless the metallic chloride formed separates from the mass of

the active substances. Thus, for example, if a solution of common

salt be mixed with a solution of magnesium sulphate, double

decomposition ensues, but not completely, because all the

substances remain in the solution. In this case the decomposition

must result in the formation of sodium sulphate and magnesium

chloride, substances which are soluble in water; nothing is

disengaged, and therefore the decomposition 2NaCl + MgSO_{4} =

MgCl_{2} + Na_{2}SO_{4} cannot proceed to the end. However, the

sodium sulphate formed in this manner may be separated by freezing

the mixture. The complete separation of the sodium sulphate will

naturally not take place, owing to a portion of the salt remaining

in the solution. Nevertheless, this kind of decomposition is made

use of for the preparation of sodium sulphate from the residues

left after the evaporation of sea-water, which contain a mixture

of magnesium sulphate and common salt. Such a mixture is found at

Stassfurt in a natural form. It might be said that this form of

double decomposition is only accomplished with a change of

temperature; but this would not be true, as may be concluded from

other analogous cases. Thus, for instance, a solution of copper

sulphate is of a blue colour, while a solution of copper chloride

is green. If we mix the two salts together the green tint is

distinctly visible, so that by this means the presence of the

copper chloride in the solution of copper sulphate is clearly

seen. If now we add a solution of common salt to a solution of

copper sulphate, a green coloration is obtained, which indicates

the formation of copper chloride. In this instance it is not

separated, but it is immediately formed on the addition of common

salt, as it should be according to Berthollet's doctrine.

The complete formation of a metallic chloride from common salt can

only occur, judging from the above, when it separates from the

sphere of action. The salts of silver are instances in point,

because the silver chloride is insoluble in water; and therefore

if we add a solution of sodium chloride to a solution of a silver

salt, silver chloride and the sodium salt of that acid which was

in the silver salt are formed.

As a proof that double decompositions like the above are actually

accomplished in the sense of Berthollet's doctrine, the fact may be cited

that common salt may be entirely decomposed by nitric acid, and nitre may

be completely decomposed by hydrochloric acid, just as they are

decomposed by sulphuric acid; but this only takes place when, in the

first instance, an excess of nitric acid is taken, and in the second

instance, an excess of hydrochloric acid, for a given quantity of the

sodium salt, and when the resultant acid passes off. If sodium chloride

be put into a porcelain evaporating basin, nitric acid added to it, and

the mixture heated, then both hydrochloric and nitric acids are expelled

by the heat. Thus the nitric acid partially acts on the sodium chloride,

but on heating, as both acids are volatile, they are both converted into

vapour; and therefore the residue will contain a mixture of a certain

quantity of the sodium chloride taken and of the sodium nitrate formed.

If a fresh quantity of nitric acid be then added, reaction will again set

in, a certain portion of hydrochloric acid is again evolved, and on

heating is expelled together with nitric acid. If this be repeated

several times, it is possible to expel all the hydrochloric acid, and to

obtain sodium nitrate only in the residue. If, on the contrary, we take

sodium nitrate and add hydrochloric acid to it in an aqueous solution, a

certain quantity of the hydrochloric acid displaces a portion of the

nitric acid, and on heating the excess of hydrochloric acid passes away

with the nitric acid formed. On repeating this process, it is possible to

displace the nitric acid with an excess of hydrochloric acid, just as it

was possible to displace the hydrochloric acid by an excess of nitric

acid. The influence of the mass of the substance in action and the

influence of volatility are here very distinctly seen. Hence it may be

affirmed that sulphuric acid does not displace hydrochloric acid because

of an especially high degree of affinity, but that this reaction is only

carried on to the end because the sulphuric acid is not volatile, whilst

the hydrochloric acid which is formed is volatile.

The preparation of hydrochloric acid in the laboratory and on a large

scale is based upon these data. In the first instance, an excess of

sulphuric acid is employed in order that the reaction may proceed easily

at a low temperature, whilst on a large scale, when it is necessary to

economise every material, equivalent quantities are taken in order to

obtain the normal salt Na_{2}SO_{4} and not the acid salt, which would

require twice as much acid. The hydrochloric acid evolved is a gas which

is very soluble in water. It is most frequently used in practice in this

state of solution under the name of _muriatic acid_.[31]

[31] The apparatus shown in fig. 46 (Chapter VI., Note 12) is generally

employed for the preparation of small quantities of hydrochloric

acid. Common salt is placed in the retort; the salt is generally

previously fused, as it otherwise froths and boils over in the

apparatus. When the apparatus is placed in order sulphuric acid

mixed with water is poured down the thistle funnel into the

retort. Strong sulphuric acid (about half as much again as the

weight of the salt) is usually taken, and it is diluted with a

small quantity of water (half) if it be desired to retard the

action, as in using strong sulphuric acid the action immediately

begins with great vigour. The mixture, at first without the aid of

heat and then at a moderate temperature (in a water-bath), evolves

hydrochloric acid. Commercial hydrochloric acid contains many

impurities; it is usually purified by distillation, the middle

portions being collected. It is purified from arsenic by adding

FeCl_{2}, distilling, and rejecting the first third of the

distillate. If free hydrochloric acid gas be required, it is

passed through a vessel containing strong sulphuric acid to dry

it, and is collected over a mercury hath.

Phosphoric anhydride absorbs hydrogen chloride (Bailey and Fowler,

1888; 2P_{2}O_{3} + 3HCl = POCl_{3} + 3HPO_{3}) at the ordinary

temperature, and therefore the gas cannot he dried by this

substance.

[Illustration: FIG. 65.--Section of a salt-cake furnace. B, pan in which

the sodium chloride and sulphuric acid are first mixed and heated. C,

muffle for the ultimate decomposition.]

In chemical works the decomposition of sodium chloride by means of

sulphuric acid is carried on on a very large scale, chiefly with a

view to the preparation of normal sodium sulphate, the hydrochloric

acid being a bye-product.[31 bis] The furnace employed is termed a

_salt cake furnace_. It is represented in fig. 65, and consists of the

following two parts: the pan B and the roaster C, or enclosed space

built up of large bricks _a_ and enveloped on all sides by the smoke

and flames from the fire grate, F. The ultimate decomposition of the

salt by the sulphuric acid is accomplished in the roaster. But the first

decomposition of sodium chloride by sulphuric acid does not require

so high a temperature as the ultimate decomposition, and is therefore

carried on in the front and cooler portion, B, whose bottom is heated by

gas flues. When the reaction in this portion ceases and the evolution of

hydrochloric acid stops, then the mass, which contains about half of the

sodium chloride still undecomposed, and the sulphuric acid in the form

of acid sodium sulphate, is removed from B and thrown into the roaster

C, where the action is completed. Normal sodium sulphate, which we

shall afterwards describe, remains in the roaster. It is employed both

directly in the manufacture of glass, and in the preparation of other

sodium compounds--for instance, in the preparation of soda ash, as will

afterwards be described. For the present we will only turn our attention

to the hydrochloric acid evolved in B and C.

[31 bis] In chemical works where sulphuric acid of 60° Baumé (22 p.c.

of water) is employed, 117 parts of sodium chloride are taken to

about 125 parts of sulphuric acid.

The hydrochloric acid gas evolved is subjected to condensation by

dissolving it in water.[32] If the apparatus in which the decomposition

is accomplished were hermetically closed, and only presented one outlet,

then the escape of the hydrochloric acid would only proceed through the

escape pipe intended for this purpose. But as it is impossible to

construct a perfectly hermetically closed furnace of this kind, it is

necessary to increase the draught by artificial means, or to oblige the

hydrochloric acid gas to pass through those arrangements in which it is

to be condensed. This is done by connecting the ends of the tubes through

which the hydrochloric acid gas escapes from the furnace with high

chimneys, where a strong draught is set up from the combustion of the

fuel. This causes a current of hydrochloric acid gas to pass through the

absorbing apparatus in a definite direction. Here it encounters a current

of water flowing in the opposite direction, by which it is absorbed. It

is not customary to cause the acid to pass through the water, but only to

bring it into contact with the surface of the water. The absorption

apparatus consists of large earthenware vessels having four orifices, two

above and two lateral ones in the wide central portion of each vessel.

The upper orifices serve for connecting the vessels together, and the

hydrochloric acid gas escaping from the furnace passes through these

tubes. The water for absorbing the acid enters at the upper, and flows

out from the lower, vessel, passing through the lateral orifices in the

vessels. The water flows from the chimney towards the furnace and it is

therefore evident that the outflowing water will be the most saturated

with acid, of which it actually contains about 20 per cent. The

absorption in these vessels is not complete. The ultimate absorption of

the hydrochloric acid is carried on in the so-called _coke towers_, which

usually consist of two adjacent chimneys. A lattice-work of bricks is

laid on the bottom of these towers, on which coke is piled up to the top

of the tower. Water, distributing itself over the coke, trickles down to

the bottom of the tower, and in so doing absorbs the hydrochloric acid

gas rising upwards.

[32] As in works which treat common salt in order to obtain sodium

sulphate, the hydrochloric acid is sometimes held to be of no

value, it might be allowed to escape with the waste furnace gases

into the atmosphere, which would greatly injure the air of the

neighbourhood and destroy all vegetation. In all countries,

therefore, there are laws forbidding the factories to proceed in

this manner, and requiring the absorption of the hydrochloric acid

by water at the works themselves, and not permitting the solution

to be run into rivers and streams, whose waters it would spoil. It

may be remarked that the absorption of hydrochloric acid presents

no particular difficulties (the absorption of sulphurous acid is

much more difficult) because hydrochloric acid has a great

affinity for water and gives a hydrate which boils above 100°.

Hence, even steam and hot water, as well as weaker solutions, can

be used for absorbing the acid. However, Warder (1888) showed that

weak solutions of composition H_{2}O + _n_HCl when boiled (the

residue will be almost HCl,8H_{2}O) evolve (not water but) a

solution of the composition H_{2}O + 445_n_^{4}HCl; for example,

on distilling HCl,10H_{2}O, HCl,23H_{2}O is first obtained in the

distillate. As the strength of the residue becomes greater, so

also does that of the distillate, and therefore in order to

completely absorb hydrochloric acid it is necessary in the end to

have recourse to water.

As in Russia the manufacture of sodium sulphate from sodium

chloride has not yet been sufficiently developed, and as

hydrochloric acid is required for many technical purposes (for

instance, for the preparation of zinc chloride, which is employed

for soaking railway sleepers), therefore salt is often treated

mainly for the manufacture of hydrochloric acid.

It will be readily understood that hydrochloric acid may be obtained

from all other metallic chlorides.[33] It is frequently formed in other

reactions, many of which we shall meet with in the further course of

this work. It is, for instance, formed by the action of water on sulphur

chloride, phosphorus chloride, antimony chloride, &c.

[33] Thus the metallic chlorides, which are decomposed to a greater or

less degree by water, correspond with feeble bases. Such are, for

example, MgCl_{2}, AlCl_{3}, SbCl_{3}, BiCl_{3}. The decomposition

of magnesium chloride (and also carnallite) by sulphuric acid

proceeds at the ordinary temperature; water decomposes MgCl_{2} to

the extent of 50 p.c. when aided by heat, and _may be employed_ as

a convenient _method for the production of hydrochloric acid_.

Hydrochloric acid is also produced by the ignition of certain

metallic chlorides in a stream of hydrogen, especially of those

metals which are easily reduced and difficultly oxidised--for

instance, silver chloride. Lead chloride, when heated to redness

in a current of steam, gives hydrochloric acid and lead oxide. The

multitude of the cases of formation of hydrochloric acid are

understood from the fact that it is a substance which is

comparatively very stable, resembling water in this respect, and

even most probably more stable than water, because, at a high

temperature and even under the action of light, chlorine

decomposes water, with the formation of hydrochloric acid. The

combination of chlorine and hydrogen also proceeds by their direct

action, as we shall afterwards describe.

_Hydrochloric acid_ is a colourless gas having a pungent suffocating

odour and an acid taste. This gas fumes in air and attracts moisture,

because it forms vapour containing a compound of hydrochloric acid and

water. Hydrochloric acid is liquefied by cold, and under a pressure of

40 atmospheres, into a colourless liquid of sp. gr. 0·908 at 0°,[34]

boiling point -35° and absolute boiling point +52°. We have already

seen (Chapter I.) that hydrochloric acid combines very energetically

_with water_, and in so doing evolves a considerable amount of heat.

The solution saturated in the cold attains a density 1·23. On heating

such a solution containing about 45 parts of acid per 100 parts, the

hydrochloric acid gas is expelled with only a slight admixture of

aqueous vapour. But it is impossible to entirely separate the whole of

the hydrochloric acid from the water by this means, as could be done

in the case of an ammoniacal solution. The temperature required for

the evolution of the gas rises and reaches 110°-111°, and after this

remains constant--that is, a solution having a constant boiling point

is obtained (as with HNO_{3}), which, however, does not (Roscoe and

Dittmar) present a constant composition under different pressures,

because the hydrate is decomposed in distillation, as is seen from the

determinations of its vapour density (Bineau). Judging from the facts

(1) that with decrease of the pressure under which the distillation

proceeds the solution of constant boiling point approaches to a

composition of 25 p.c. of hydrochloric acid,[35] (2) that by passing

a stream of dry air through a solution of hydrochloric acid there is

obtained in the residue a solution which also approaches to 25 p.c. of

acid, and more nearly as the temperature falls,[36] (3) that many of the

properties of solutions of hydrochloric acid vary distinctly according

as they contain more or less than 25 p.c. of hydrochloric acid (for

instance, antimonious sulphide gives hydrogen sulphide with a stronger

acid, but is not acted on by a weaker solution, also a stronger solution

fumes in the air, &c.), and (4) that the composition HCl,6H_{2}O

corresponds with 25·26 p.c. HCl--judging from all these data, and also

from the loss of tension which occurs in the combination of hydrochloric

acid with water, it may be said that they form a _definite hydrate_ of

the composition HCl,6H_{2}O. Besides this hydrate there exists also a

crystallo-hydrate, HCl,2H_{2}O,[37] which is formed by the absorption of

hydrochloric acid by a saturated solution at a temperature of -23°. It

crystallises and melts at -18°.[38]

[34] According to Ansdell (1880) the sp. gr. of liquid hydrochloric

acid at 0° = 0·908, at 11·67° = 0·854, at 22·7° = 0·808, at 33° =

0·748. Hence it is seen that the expansion of this liquid is

greater than that of gases (Chapter II., Note 34).

[35] According to Roscoe and Dittmar at a pressure of three atmospheres

the solution of constant boiling point contains 18 p.c. of

hydrogen chloride, and at a pressure of one-tenth atmosphere 23

p.c. The percentage is intermediate at medium pressures.

[36] At 0° 25 p.c., at 100° 20·7 p.c.; Roscoe and Dittmar.

[37] This crystallo-hydrate (obtained by Pierre and Puchot, and

investigated by Roozeboom) is analogous to NaCl,2H_{2}O. The

crystals HCl,2H_{2}O at -22° have a specific gravity 1·46; the

vapour tension (under dissociation) of the solution having a

composition HCl,2H_{2}O at -24° = 760, at -19° = 1,010, at -18° =

1,057, at -17° = 1,112 mm. of mercury. In a solid state the

crystallo-hydrate at -17·7° has the same tension, whilst at lower

temperatures it is much less: at -24° about 150, at -19° about 580

mm. A mixture of fuming hydrochloric acid with snow reduces the

temperature to -38°. If another equivalent of water be added to the

hydrate HCl,2H_{2}O at -18°, the temperature of solidification

falls to -25°, and the hydrate HCl,3H_{2}O is formed (Pickering,

1893).

[38] According to Roscoe at 0° one _hundred_ grams of water at a

pressure _p_ (in millimetres of mercury) dissolves--

_p_ = 100 200 300 500 700 1,000

Grams HCl 65·7 70·7 73·8 78·2 81·7 85·6

At a pressure of 760 millimetres and temperature _t_, one

_hundred_ grams of water dissolves

_t_ = 0 8° 16° 24° 40° 60°

Grams HCl 82·5 78·3 74·2 70·0 63·3 56·1

Roozeboom (1886) showed that at _t_° solutions containing _c_

grams of hydrogen chloride per 100 grams of water may (with the

variation of the pressure _p_) be formed together with the

crystallo-hydrate HCl,2H_{2}O:

_t_ = -28°·8 -21° -19° -18° -17°·7

_c_ = 84·2 86·8 92·6 98·4 101·4

_p_ = -- 334 580 900 1,073 mm.

The last combination answers to the melted crystallo-hydrate

HCl,2H_{2}O, which splits up at temperatures above -17°·7, and at

a constant atmospheric pressure when there are no crystals--

_t_ = -24° -21° -18° -10° 0°

_c_ = 101·2 98·3 95·7 89·8 84·2

From these data it is seen that the hydrate HCl,2H_{2}O can exist

in a liquid state, which is not the case for the hydrates of

carbonic and sulphurous anhydrides, chlorine, &c.

According to Marignac, the specific heat _c_ of a solution HCl +

_m_H_{2}O (at about 30°, taking the specific heat of water = 1) is

given by the expression--

C(36·5 + _m_18) = 18_m_ - 28·39 + 140/_m_ - 268/_m_^2

if _m_ be not less than 6·25. For example, for HCl + 25H_{2}O,

C = 0·877.

According to Thomsen's data, the amount of heat _Q_, expressed in

thousands of calories, evolved in the solution of 36·5 grams of

gaseous hydrochloric acid in _m_H_{2}O or 18_m_ grams of water is

equal to--

_m_ = 2 4 10 50 400

_Q_ = 11·4 14·3 16·2 17·1 17·3

In these quantities the latent heat of liquefaction is included,

which must be taken as 5-9 thousand calories per molecular

quantity of hydrogen chloride.

The researches of Scheffer (1888) on the rate of diffusion (in

water) of solutions of hydrochloric acid show that the coefficient

of diffusion _k_ decreases with the amount of water _n_, if the

composition of the solution is HCl,_n_H_{2}O at 0°:--

_n_ = 5 6·9 9·8 14 27·1 129·5

_k_ = 2·31 2·08 1·86 1·67 1·52 1·39

It also appears that strong solutions diffuse more rapidly into

dilute solutions than into water.

The mean specific gravities at 15°, taking water at its maximum density

(4°) as 10,000, for solutions containing _p_ per cent. of hydrogen

chloride are--

_p_ _S_ _p_ _S_

5 10,242 25 11,266

10 10,490 30 11,522

15 10,744 35 11,773

20 11,001 40 11,997

The formula _S_ = 9,991·6 + 49·43_p_ + 0·0571_p_^2, up to _p_ =

25·26, which answers to the hydrate HCl,6H_{2}O mentioned above,

gives the specific gravity. Above this percentage _S_ = 9,785·1 +

65·10_p_-0·240_p_^2. The rise of specific gravity with an increase

of percentage (or the differential _ds/dp_) reaches a maximum at

about 25 p.c.[39] The intermediate solution, HCl,6H_{2}O, is further

distinguished by the fact that the variation of the specific gravity

with the variation of temperature is a constant quantity, so that the

specific gravity of this solution is equal to 11,352·7(1-0·000447_t_),

where 0·000447 is the coefficient of expansion of the solution.[40]

In the case of more dilute solutions, as with water, the specific

gravity per 1° (or the differential _ds_/_dt_) rises with a rise of

temperature.[41]

_p_ = 0 5 10 15 20

_S__{0} - _S__{15} = 7·2 23 38 52 64

_S__{15} - _S__{30} = 34·1 42 50 59 67

Whilst for solutions which contain a greater proportion of hydrogen

chloride than HCl,6H_{2}O, these coefficients _decrease_ with a

rise of temperature; for instance, for 30 p.c. of hydrogen chloride

_S__{0}-_S__{15} = 88 and _S__{15}-_S__{30} = 87 (according to

Marignac's data). In the case of HCl,6H_{2}O these differences are

constant, and equal 76.

[39] If it be admitted that the maximum of the differential corresponds

with HCl,6H_{2}O, then it might be thought that the specific

gravity is expressed by a parabola of the third order; but such an

admission does not give expressions in accordance with fact. This

is all more fully considered in my work mentioned in Chapter I.,

Note 19.

[40] As in water, the coefficient of expansion (or the quantity _k_ in

the expression S_{_t_} = S_{_0_}-_k_S_{_0_}_t_, or V_t_ =

1/(1-_kt_)) attains a magnitude 0·000447 at about 48°, it might be

thought that at 48° all solutions of hydrochloric acid would have

the same coefficient of expansion, but in reality this is not the

case. At low and at the ordinary temperatures the coefficient of

expansion of aqueous solutions is greater than that of water, and

increases with the amount of substance dissolved.

[41] The figures cited above may serve for the direct determination of

that variation of the specific gravity of solutions of

hydrochloric acid with the temperature. Thus, knowing that at 15°

the specific gravity of a 10 p.c. solution of hydrochloric acid =

10,492, we find that at _t_° it = 10,530-_t_(2·13 + 0·027_t_).

Whence also may be found the coefficient of expansion (Note 40).

Thus the formation of two definite hydrates, HCl,2H_{2}O and HCl,6H_{2}O,

between hydrochloric acid and water may be accepted upon the basis of

many facts. But both of them, if they occur in a liquid state, dissociate

with great facility into hydrogen chloride and water, and are completely

decomposed when distilled.

All solutions of hydrochloric acid present the properties of an

energetic acid. They not only transform blue vegetable colouring matter

into red, and disengage carbonic acid gas from carbonates, &c., but

they also entirely saturate bases, even such energetic ones as potash,

lime, &c. In a dry state, however, hydrochloric acid does not alter

vegetable dyes, and does not effect many double decompositions which

easily take place in the presence of water. This is explained by the

fact that the gaso-elastic state of the hydrochloric acid prevents its

entering into reaction. However, incandescent iron, zinc, sodium, &c.,

act on gaseous hydrochloric acid, displacing the hydrogen and leaving

half a volume of hydrogen for each volume of hydrochloric acid gas; this

reaction may serve for determining the composition of hydrochloric acid.

Combined with water hydrochloric acid acts as an acid much resembling

nitric acid[42] in its energy and in many of its reactions; however, the

latter contains oxygen, which is disengaged with great ease, and so very

frequently acts as an oxidiser, which hydrochloric acid is not capable

of doing. The majority of metals (even those which do not displace the

H from H_{2}SO_{4}, but which, like copper, decompose it to the limit

of SO_{2}) displace the hydrogen from hydrochloric acid. Thus hydrogen

is disengaged by the action of zinc, and even of copper and tin.[42 bis]

Only a few metals withstand its action; for example, gold and

platinum. Lead in compact masses is only acted on feebly, because the

lead chloride formed is insoluble and prevents the further action of the

acid on the metal. The same is to be remarked with respect to the feeble

action of hydrochloric acid on mercury and silver, because the compounds

of these metals, AgCl and HgCl, are insoluble in water. Metallic

chlorides are not only formed by the action of hydrochloric acid on the

metals, but also by many other methods; for instance, by the action of

hydrochloric acid on the carbonates, oxides, and hydroxides, and also

by the action of chlorine on metals and certain of their compounds.

Metallic chlorides have a composition MCl; for example, NaCl, KCl,

AgCl, HgCl, if the metal replaces hydrogen equivalent for equivalent,

or, as it is said, if it be monatomic or univalent. In the case of

bivalent metals, they have a composition MCl_{2}; for example, CaCl_{2},

CuCl_{2}, PbCl_{2}, HgCl_{2}, FeCl_{2}, MnCl_{2}. The composition of

the haloid salts of other metals presents a further variation; for

example, AlCl_{3}, PtCl_{4}, &c. Many metals, for instance Fe, give

several degrees of combination with chlorine (FeCl_{2}, FeCl_{3}) as

with hydrogen. In their composition the metallic chlorides differ

from the corresponding oxides, in that the O is replaced by Cl_{2},

as should follow from the law of substitution, because oxygen gives

OH_{2}, and is consequently bivalent, whilst chlorine forms HCl, and is

therefore univalent. So, for instance, ferrous oxide, FeO, corresponds

with ferrous chloride, FeCl_{2}, and the oxide Fe_{2}O_{3} with ferric

chloride, which is also seen from the origin of these compounds, for

FeCl_{2} is obtained by the action of hydrochloric acid on ferrous

oxide or carbonate and FeCl_{3} by its action on ferric oxide. In a

word, all the typical properties of acids are shown by hydrochloric

acid, and all the typical properties of salts in the metallic chlorides

derived from it. Acids and salts composed like HCl and M_{n}Cl_{2m}

without any oxygen bear the name of haloid salts; for instance, HCl is

a haloid acid, NaCl a haloid salt, chlorine a halogen. The capacity

of hydrochloric acid to give, by its action on bases, MO, a metallic

chloride, MCl_{2}, and water, is limited at high temperatures by the

reverse reaction MCl_{2} + H_{2}O = MO + 2HCl, and the more pronounced

are the basic properties of MO the feebler is the reverse action, while

for feebler bases such as Al_{2}O_{3}, MgO, &c., this reverse reaction

proceeds with ease. Metallic chlorides corresponding with the peroxides

either do not exist, or are easily decomposed with the disengagement

of chlorine. Thus there is no compound BaCl_{4} corresponding with

the peroxide BaO_{2}. Metallic chlorides having the general aspect of

salts, like their representative sodium chloride, are, as a rule, easily

fusible, more so than the oxides (for instance, CaO is infusible at a

furnace heat, whilst CaCl_{2} is easily fused) and many other salts.

Under the action of heat many chlorides are more stable than the oxides,

some can even be converted into vapour; thus corrosive sublimate,

HgCl_{2}, is particularly volatile, whilst the oxide HgO decomposes at

a red heat. Silver chloride, AgCl, is fusible and is decomposed with

difficulty, whilst Ag_{2}O is easily decomposed. The majority of the

metallic chlorides are soluble in water, but silver chloride, cuprous

chloride, mercurous chloride, and lead chloride are sparingly soluble in

water, and are therefore easily obtained as precipitates when a solution

of the salts of these metals is mixed with a solution of any chloride

or even with hydrochloric acid. The metal contained in a haloid salt

may often be replaced by another metal, or even by hydrogen, just as is

the case with a metal in an oxide. Thus copper displaces mercury from

a solution of mercuric chloride, HgCl_{2} + Cu = CuCl_{2} + Hg, and

hydrogen at a red heat displaces silver from silver chloride, 2AgCl +

H_{2} = Ag_{2} + 2HCl. These, and a whole series of similar reactions,

form the typical methods of double saline decompositions. The measure of

decomposition and the conditions under which reactions of double saline

decompositions proceed in one or in the other direction are determined

by the properties of the compounds which take part in the reaction, and

of those capable of formation at the temperature, &c., as was shown in

the preceding portions of this chapter, and as will be frequently found

hereafter.

[42] Thus, for instance, with feeble bases they evolve in dilute

solutions (Chapter III., Note 53) almost equal amounts of heat;

their relation to sulphuric acid is quite identical. They both

form fuming solutions as well as hydrates; they both form

solutions of constant boiling point.

[42 bis] Pybalkin (1891) found that copper begins to disengage hydrogen

at 100°, and that chloride of copper begins to give up its

chlorine to hydrogen gas at 230°; for silver these temperatures

are 117° and 260°--that is, there is less difference between them.

If hydrochloric acid enters into double decomposition with basic oxides

and their hydrates, this is only due to its acid properties; and for the

same reason it rarely enters into double decomposition with acids and

acid anhydrides. Sometimes, however, it combines with the latter, as, for

instance, with the anhydride of sulphuric acid, forming the compound

SO_{3}HCl; and in other cases it acts on acids, giving up its hydrogen to

their oxygen and forming chlorine, as will be seen in the following

chapter.

Hydrochloric acid, as may already be concluded from the composition of

its molecule, belongs to the monobasic acids, and does not, therefore,

give true acid salts (like HNaSO_{4} or HNaCO_{3}); nevertheless

many metallic chlorides, formed from powerful bases, are capable of

_combining with hydrochloric acid_, just as they combine with water,

or with ammonia, or as they give double salts. Compounds have long

been known of hydrochloric acid with auric, platinic, and antimonious

chlorides, and other similar metallic chlorides corresponding with

very feeble bases. But Berthelot, Engel, and others have shown that

the capacity of HCl for combining with M_{_n_}Cl_{_m_} is much more

frequently encountered than was previously supposed. Thus, for

instance, dry hydrochloric acid when passed into a solution of zinc

chloride (containing an excess of the salt) gives in the cold (0°)

a compound HCl,ZnCl_{2},2H_{2}O, and at the ordinary temperature

HCl,2ZnCl_{2},2H_{2}O, just as it is able at low temperatures to

form the crystallo-hydrate ZnCl_{2},3H_{2}O (Engel, 1886). Similar

compounds are obtained with CdCl_{2},CuCl_{2}, HgCl_{2},Fe_{2}Cl_{6},

&c. (Berthelot, Ditte, Cheltzoff, Lachinoff, and others). These

compounds with hydrochloric acid are generally more soluble in water

than the metallic chlorides themselves, so that whilst hydrochloric

acid decreases the solubility of M_{_n_}Cl_{_m_}, corresponding with

energetic bases (for instance, sodium or barium chlorides), it increases

the solubility of the metallic chlorides corresponding with feeble bases

(cadmium chloride, ferric chloride, &c.) Silver chloride, which is

insoluble in water, is soluble in hydrochloric acid. Hydrochloric acid

also combines with certain unsaturated hydrocarbons (for instance, with

turpentine, C_{10}H_{16},2HCl) and their derivatives. _Sal-ammoniac_,

or ammonia hydrochloride, NH_{4}Cl = NH_{3},HCl, also belongs to

this class of compounds.[43] If hydrogen chloride gas be mixed with

ammonia gas a solid compound consisting of equal volumes of each is

immediately formed. The same compound is obtained on mixing solutions of

the two gases. It is also produced by the action of hydrochloric acid

on ammonium carbonate. Sal-ammoniac is usually prepared, in practice,

by the last method.[44] The specific gravity of sal-ammoniac is 1·55.

We have already seen (Chapter VI.) that sal-ammoniac, like all other

ammonium salts, easily decomposes; for instance, by volatilisation with

alkalis, and even partially when its solution is boiled. The other

properties and reactions of sal-ammoniac, especially in solution, fully

recall those already mentioned in speaking of sodium chloride. Thus, for

instance, with silver nitrate it gives a precipitate of silver chloride;

with sulphuric acid it gives hydrochloric acid and ammonium sulphate,

and it forms double salts with certain metallic chlorides and other

salts.[45]

[43] When an unsaturated hydrocarbon, or, in general, an unsaturated

compound, assimilates to itself the molecules Cl_{2}, HCl, SO_{3},

H_{2}SO_{4}, &c., the cause of the reaction is most simple. As

nitrogen, besides the type NX_{3} to which NH_{3}, belongs, gives

compounds of the type NX_{5}--for example, NO_{2}(OH)--the

formation of the salts of ammonium should be understood in this

way. NH_{3} gives NH_{4}Cl because NX_{3} is capable of giving

NX_{5}. But as saturated compounds--for instance, SO_{3},H_{2}O,

NaCl, &c.--are also capable of combination even between

themselves, it is impossible to deny the capacity of HCl also for

combination. SO_{3} combines with H_{2}O, and also with HCl and

the unsaturated hydrocarbons. It is impossible to recognise the

distinction formerly sought to be established between atomic and

molecular compounds, and regarding, for instance, PCl_{3} as an

atomic compound and PCl_{5} as a molecular one, only because it

easily splits up into molecules PCl_{3} and Cl_{2}.

[44] Sal-ammoniac is prepared from ammonium carbonate, obtained in the

dry distillation of nitrogenous substances (Chapter VI.), by

saturating the resultant solution with hydrochloric acid. A

solution of sal-ammoniac is thus produced, which is evaporated,

and in the residue a mass is obtained containing a mixture of

various other, especially tarry, products of dry distillation. The

sal-ammoniac is generally purified by sublimation. For this

purpose iron vessels covered with hemispherical metallic covers

are employed, or else simply clay crucibles covered by other

crucibles. The upper portion, or head, of the apparatus of this

kind will have a lower temperature than the lower portion, which

is under the direct action of the flame. The sal-ammoniac

volatilises when heated, and settles on the cooler portion of the

apparatus. It is thus freed from many impurities, and is obtained

as a crystalline crust, generally several centimetres thick, in

which form it is commonly sold. The solubility of sal-ammoniac

rises rapidly with the temperature: at 0°, 100 parts of water

dissolve about 28 parts of NH_{4}Cl, at 50° about 50 parts, and at

the ordinary temperature about 35 parts. This is sometimes taken

advantage of for separating NH_{4}Cl from solutions of other

salts.

[45] The solubility of sal-ammoniac in 100 parts of water (according to

Alluard) is--

0° 10° 20° 30° 40° 60° 80° 100° 110°

28·40 32·48 37·28 41·72 46 55 64 73 77

A saturated solution boils at 115°·8. The specific gravity at

15°/4° of solutions of sal-ammoniac (water 4° = 10,000) =

9,991·6-31·26_p_-0·085_p_^2, where _p_ is the amount by weight of

ammonium chloride in 100 parts of solution. With the majority of

salts the differential _ds_/_dp_ increases, but here it decreases

with the increase of _p_. For (unlike the sodium and potassium

salts) a solution of the alkali _plus_ a solution of acid occupy a

greater volume than that of the resultant ammonium salt. In the

solution of _solid_ ammonium chloride a contraction, and not

expansion, generally takes place. It may further be remarked that

solutions of sal-ammoniac have an acid reaction even when prepared

from the salt remaining after prolonged washing of the sublimed

salt with water (A. Stcherbakoff).

CHAPTER XI

THE HALOGENS: CHLORINE, BROMINE, IODINE, AND FLUORINE

Although hydrochloric acid, like water, is one of the most stable

substances, it is nevertheless decomposed not only by the action of

a galvanic current,[1] but also by a high temperature. Sainte-Claire

Deville showed that decomposition already occurs at 1,300°, because a

cold tube (as with CO, Chapter IX.) covered with an amalgam of silver

absorbs chlorine from hydrochloric acid in a red-hot tube, and the

escaping gas contains hydrogen. V. Meyer and Langer (1885) observed

the decomposition of hydrochloric acid at 1,690° in a platinum vessel;

the decomposition in this instance was proved not only from the fact

that hydrogen diffused through the platinum (p. 142), owing to which

the volume was diminished, but also from chlorine being obtained in the

residue (the hydrogen chloride was mixed with nitrogen), which liberated

iodine from potassium iodide.[2] The usual method for the preparation

of chlorine consists in the abstraction of the hydrogen by oxidising

agents.[2 bis]

[1] The decomposition of fused sodium chloride by an electric current

has been proposed in America and Russia (N. N. Beketoff) as a means

for the preparation of chlorine and sodium. A strong solution of

hydrochloric acid is decomposed into equal volumes of chlorine and

hydrogen by the action of an electric current. If sodium chloride

and lead be melted in a crucible, the former being connected with

the cathode and a carbon anode immersed in the lead, then the lead

dissolves sodium and chlorine is disengaged as gas. This

electrolytic method has not yet been practised on a large scale,

probably because gaseous chlorine has not many applications, and

because of the difficulty there is in dealing with it.

[2] To obtain so high a temperature (at which the best kinds of

porcelain soften) Langer and Meyer employed the dense graphitoidal

carbon from gas retorts, and a powerful blast. They determined the

temperature by the alteration of the volume of nitrogen in the

platinum vessel, for this gas does not permeate through platinum,

and is unaltered by heat.

[2 bis] The acid properties of hydrochloric acid were known when

Lavoisier pointed out the formation of acids by the combination of

water with the oxides of the non-metals, and therefore there was

reason for thinking that hydrochloric acid was formed by the

combination of water with the oxide of some element. Hence when

Scheele obtained chlorine by the action of hydrochloric acid on

manganese peroxide he considered it as the acid contained in common

salt. When it became known that chlorine gives hydrochloric acid

with hydrogen, Lavoisier and Berthollet supposed it to be a

compound with oxygen of an anhydride contained in hydrochloric

acid. They supposed that hydrochloric acid contained water and the

oxide of a particular radicle, and that chlorine was a higher

degree of oxidation of this radicle _muvias_ (from the Latin neme

of hydrochloric acid, _acidum muriaticum_). It was only in 1811

that Gay-Lussac and Thénard in France and Davy in England arrived

at the conclusion that the substance obtained by Scheele does not

contain oxygen, nor under any conditions give water with hydrogen,

and that there is no water in hydrochloric acid gas, and therefore

concluded that chlorine is an elementary substance. They named it

'chlorine' from the Greek word [Greek: chlôros], signifying a green

colour, because of the peculiar colour by which this gas is

characterised.

An aqueous solution of hydrochloric acid is generally employed for t he

evolution of chlorine. The hydrogen has to be abstracted from the

hydrochloric acid. This is accomplished by nearly all oxidising

substances, and especially by those which are able to evolve oxygen at a

red heat (besides bases, such as mercury and silver oxides, which are

able to give salts with hydrogen chloride); for example, manganese

peroxide, potassium chlorate, chromic acid, &c. The decomposition

essentially consists in the oxygen of the oxidising substance displacing

the chlorine from 2HCl, forming water, H_{2}O, and setting the chlorine

free, 2HCl + O (disengaged by the oxidising substances) = H_{2}O +

Cl_{2}. Even nitric acid partially produces a like reaction; but as we

shall afterwards see its action is more complicated, and it is therefore

not suitable for the preparation of pure chlorine.[3] But other oxidising

substances which do not give any other volatile products with

hydrochloric acid may be employed for the preparation of chlorine. Among

these may be mentioned: potassium chlorate, acid potassium chromate,

sodium manganate, manganese peroxide, &c. Manganese peroxide is commonly

employed in the laboratory, and on a large scale, for the preparation of

chlorine. The chemical process in this case may be represented as

follows: an exchange takes place between 4HCl and MnO_{2}, in which the

manganese takes the place of the four atoms of hydrogen, or the chlorine

and oxygen exchange places--that is, MnCl_{4} and 2H_{2}O are produced.

The chlorine compound, MnCl_{4}, obtained is very unstable; it splits up

into chlorine, which as a gas passes from the sphere of action, and a

lower compound containing less chlorine than the substance first formed,

which remains in the apparatus in which the mixture is heated, MnCl_{4} =

MnCl_{2} + Cl_{2}.[3 bis] The action of hydrochloric acid requires a

temperature of about 100°. In the laboratory the _preparation of

chlorine_ is carried on in flasks, heated over a water-bath, by acting on

manganese peroxide with hydrochloric acid or a mixture of common salt

and sulphuric acid[4] and washing the gas with water to remove

hydrochloric acid.[5] Chlorine cannot be collected over mercury, because

it combines with it as with many other metals, and it is soluble in

water; however, it is but slightly soluble in hot water or brine. Owing

to its great weight, chlorine may be directly collected in a dry vessel

by carrying the gas-conducting tube down to the bottom of the vessel. The

chlorine will lie in a heavy layer at the bottom of the vessel, displace

the air, and the extent to which it fills the vessel may be followed by

its colour.[6]

[3] However, nitric acid has been proposed as a means for obtaining

chlorine, but by methods which have the drawback of being very

complicated

[3 bis] This representation of the process of the reaction is most

natural. However, this decomposition is generally represented as if

chlorine gave only one degree of combination with manganese,

MnCl_{2}, and therefore directly reacts in the following

manner--MnO_{2} + 4HCl = MnCl_{2} + 2H_{2}O + Cl_{2}, in which case

it is supposed that manganese peroxide, MnO_{2}, breaks up, as it

were, into manganous oxide, MnO and oxygen, both of which react

with hydrochloric acid, the manganous oxide acting upon HCl as a

base, giving MnCl_{2} and at the same time 2HCl + O = H_{2}O +

Cl_{2}. In reality, a mixture of oxygen and hydrochloric acid does

give chlorine at a red heat, and this reaction may also take place

at the moment of its evolution in this case.

All the oxides of manganese (Mn_{2}O_{3}, MnO_{2}, MnO_{3},

Mn_{2}O_{7}), with the exception of manganous oxide, MnO, disengage

chlorine from hydrochloric acid, because manganous chloride,

MnCl_{2}, is the only compound of chlorine and manganese which

exists as a stable compound, all the higher chlorides of manganese

being unstable and evolving chlorine. Hence we here take note of

two separate changes: (1) an exchange between oxygen and chlorine,

and (2) the instability of the higher chlorine compounds. As

(according to the law of substitution) in the substitution of

oxygen by chlorine, Cl_{2} takes the place of O, the chlorine

compounds will contain more atoms than the corresponding oxygen

compounds. It is not surprising, therefore, that certain of the

chlorine compounds corresponding with oxygen compounds do not

exist, or if they are formed are very unstable. And furthermore, an

atom of chlorine is heavier than an atom of oxygen, and therefore a

given element would have to retain a large mass of chlorine if in

the higher oxides the oxygen were replaced by chlorine. For this

reason equivalent compounds of chlorine do not exist for all oxygen

compounds. Many of the former are immediately decomposed, when

formed, with the evolution of chlorine. From this it is evident

that there should exist such chlorine compounds as would evolve

chlorine as peroxides evolve oxygen, and indeed a large number of

such compounds are known. Amongst them may be mentioned antimony

pentachloride, SbCl_{5}, which splits up into chlorine and antimony

trichloride when heated. Cupric chloride, corresponding with copper

oxide, and having a composition CuCl_{2}, similar to CuO, when

heated parts with half its chlorine, just as barium peroxide

evolves half its oxygen. This method may even be taken advantage of

for the preparation of chlorine and cuprous chloride, CuCl. The

latter attracts oxygen from the atmosphere, and in so doing is

converted from a colourless substance into a green compound whose

composition is Cu_{2}Cl_{2}O. With hydrochloric acid this substance

gives cupric chloride (Cu_{2}Cl_{2}O + 2HCl = H_{2}O + 2CuCl_{2}),

which has only to be dried and heated in order again to obtain

chlorine. Thus, in solution, and at the ordinary temperature, the

compound CuCl_{2} is stable, but when heated it splits up. On this

property is founded Deacon's process for the preparation of

chlorine from hydrochloric acid with the aid of air and copper

salts, by passing a mixture of air and hydrochloric acid at about

440° over bricks saturated with a solution of a copper salt (a

mixture of solutions of CuSO_{4} and Na_{2}SO_{4}). CuCl_{2} is

then formed by the double decomposition of the salt of copper and

the hydrochloric acid; the CuCl_{2} liberates chlorine, and the

CuCl forms Cu_{2}Cl_{2}O with the oxygen of the air, which again

gives CuCl_{2} with 2HCl, and so on.

Magnesium chloride, which is obtained from sea-water, carnallite,

&c., may serve not only as a means for the preparation of

hydrochloric acid, but also of chlorine, because its basic salt

(magnesium oxychloride) when heated in the air gives magnesium

oxide and chlorine (Weldon-Pechiney's process, 1888). Chlorine is

now prepared on a large scale by this method. Several new methods

based upon this reaction have been proposed for procuring chlorine

from the bye-products of other chemical processes. Thus, Lyte and

Tattars (1891) obtained up to 67 p.c. of chlorine from CaCl_{2} in

this manner. A solution of CaCl_{2}, containing a certain amount of

common salt, is evaporated and oxide of magnesium added to it. When

the solution attains a density of 1·2445 (at 15°), it is treated

with carbonic acid, which precipitates carbonate of calcium, while

chloride of magnesium remains in solution. After adding ammonium

chloride, the solution is evaporated to dryness and the double

chloride of magnesium and ammonium formed is ignited, which drives

off the chloride of ammonium. The chloride of magnesium which

remains behind is used in the Weldon-Pechiney process. The De

Wilde-Reychler (1892) process for the manufacture of chlorine

consists in passing alternate currents of hot air and hydrochloric

acid gas through a cylinder containing a mixture of the chlorides

of magnesium and manganese. A certain amount of sulphate of

magnesium which does not participate in any way in the reaction, is

added to the mixture to prevent its fusing. The reactions may be

expressed by the following equations: (1) 3MgCl_{2} + 3MnCl_{2} +

8O = Mg_{3}Mn_{3}O_{8} + 12Cl; (2) Mg_{3}Mn_{3}O_{8} + 16HCl =

3MgCl_{2} + 3MnCl_{2} + 8H_{2}O + 4Cl. As nitric acid is able to

take up the hydrogen from hydrochloric acid, a heated mixture of

these acids is also employed for the preparation of chlorine. The

resultant mixture of chlorine and lower oxides of nitrogen is mixed

with air and steam which regenerates the HNO_{3}, while the

chlorine remains as a gas together with nitrogen, in which form it

is quite capable of bleaching, forming chloride of lime, &c.

Besides these, Solvay and Mond's methods of preparing chlorine must

be mentioned. The first is based upon the reaction CaCl_{2} +

SiO_{2} + O(air) = CaOSiO_{2} + Cl_{2}, the second on the action of

the oxygen of the air (heated) upon MgCl_{2} (and certain similar

chlorides) MgCl_{2} + O = MgO + Cl_{2} The remaining MgO is treated

with sal-ammoniac to re-form MgCl_{2} (MgO + 2NH_{4}Cl = MgCl_{2} +

H_{2}O + 2NH_{3}) and the resultant NH_{3} again converted into

sal-ammoniac, so that hydrochloric acid is the only substance

consumed. The latter processes have not yet found much application.

[4] The following proportions are accordingly taken by weight: 5 parts

of powdered manganese peroxide, 11 parts of salt (best fused, to

prevent its frothing), and 14 parts of sulphuric acid previously

mixed with an equal volume of water. The mixture is heated in a

salt bath, so as to obtain a temperature above 100°. The corks in

the apparatus must be soaked in paraffin (otherwise they are

corroded by the chlorine), and black india-rubber tubing smeared

with vaseline must be used, and not vulcanised rubber (which

contains sulphur, and becomes brittle under the action of the

chlorine).

The reaction which proceeds may be expressed thus: MnO_{2} + 2NaCl

+ 2H_{2}SO_{4} = MnSO_{4} + Na_{2}SO_{4} + 2H_{2}O + Cl_{2}. The

method of preparation of Cl_{2} from manganese peroxide and

hydrochloric acid was discovered by Scheele, and from sodium

chloride by Berthollet.

[5] The reaction of hydrochloric acid upon bleaching powder gives

chlorine without the aid of heat, CaCl_{2}O_{2} + 4HCl = CaCl_{2} +

2H_{2}O + 2Cl_{2} and is therefore also used for the preparation of

chlorine. This reaction is very violent if all the acid be added at

once; it should be poured in drop by drop (Mermé, Kämmerer). C.

Winkler proposed to mix bleaching powder with one quarter of burnt

and powdered gypsum, and having damped the mixture with water, to

press and cut it up into cubes and dry at the ordinary temperature.

These cubes can be used for the preparation of chlorine in the same

apparatus as that used for the evolution of hydrogen and carbonic

anhydride--the disengagement of the chlorine proceeds uniformly.

A mixture of potassium dichromate and hydrochloric acid evolves

chlorine perfectly free from oxygen (V. Meyer and Langer).

[6] [Illustration: FIG. 66.--Clay retort for the preparation of chlorine

on a large scale.]

Chlorine is manufactured on a _large scale_ from manganese peroxide

and hydrochloric acid. It is most conveniently prepared in the

apparatus shown in fig. 66, which consists of a three-necked

earthenware vessel whose central orifice is the largest. A clay or

lead funnel, furnished with a number of orifices, is placed in the

central wide neck of the vessel. Roughly-ground lumps of natural

manganese peroxide are placed in the funnel, which is then closed

by the cover N, and luted with clay. One orifice is closed by a

clay stopper, and is used for the introduction of the hydrochloric

acid and withdrawal of the residues. The chlorine disengaged passes

along a leaden gas-conducting tube placed in the other orifice. A

row of these vessels is surrounded by a water-bath to ensure their

being uniformly heated. Manganese chloride is found in the residue.

In Weldon's process lime is added to the acid solution of manganese

chloride. A double decomposition takes place, resulting in the

formation of manganous hydroxide and calcium chloride. When the

insoluble manganous hydroxide has settled, a further excess of milk

of lime is added (to make a mixture 2Mn(OH)_{2} + CaO +

_x_CaCl_{2}, which is found to be the best proportion, judging from

experiment), and then air is forced through the mixture. The

hydroxide is thus converted from a colourless to a brown substance,

containing peroxide, MnO_{2}, and oxide of manganese, Mn_{2}O_{3}.

This is due to the manganous oxide absorbing oxygen from the air.

Under the action of hydrochloric acid this mixture evolves

chlorine, because of all the compounds of chlorine and manganese

the chloride MnCl_{2} is the only one which is stable (_see_ Note

3). Thus one and the same mass of manganese may be repeatedly used

for the preparation of chlorine. The same result is attained in

other ways. If manganous oxide be subjected to the action of oxides

of nitrogen and air (Coleman's process), then manganese nitrate is

formed, which at a red heat gives oxides of nitrogen (which are

again used in the process) and manganese peroxide, which is thus

renewed for the fresh evolution of chlorine.

Chlorine is a _gas_ of a yellowish green colour, and has a very

suffocating and characteristic odour. On lowering the temperature to

-50° or increasing the pressure to six atmospheres (at 0°) chlorine

condenses[7] into a liquid which has a yellowish-green colour, a density

of 1·3, and boils at -34°. The density and atomic weight of chlorine is

35·5 times greater than that of hydrogen, hence the molecule contains

Cl_{2}[8]. At 0° one volume of water dissolves about 1-1/2 volume of

chlorine, at 10° about 3 volumes, at 50° again 1-1/2 volume.[9] Such

a solution of chlorine is termed 'chlorine water;' and is employed in

a diluted form in medicine and as a laboratory reagent. It is prepared

by passing chlorine through a series of Woulfe's bottles or into an

inverted retort filled with water. Under the action of light, chlorine

water gives oxygen and hydrochloric acid. At 0° a saturated solution of

chlorine yields a crystallo-hydrate, Cl_{2},8H_{2}O, which easily splits

up into chlorine and water when heated, so that if it be sealed up in a

tube and heated to 35°, two layers of liquid are formed--a lower stratum

of chlorine containing a small quantity of water, and an upper stratum

of water containing a small quantity of chlorine.[10]

[7] Davy and Faraday liquefied chlorine in 1823 by heating the

crystallo-hydrate Cl_{2}8H_{2}O in a bent tube (as with NH_{3}),

surrounded by warm water, while the other end of the tube was

immersed in a freezing mixture. Meselan condensed chlorine in

freshly-burnt charcoal (placed in a glass tube), which when cold

absorbs an equal weight of chlorine. The tube was then fused up,

the bent end cooled, and the charcoal heated, by which means the

chlorine was expelled from the charcoal, and the pressure

increased.

[8] Judging from Ludwig's observations (1868), and from the fact that

the coefficient of expansion of gases increases with their

molecular weight (Chapter II., Note 26, for hydrogen = 0·367,

carbonic anhydride = 0·373, hydrogen bromide = 0·386), it might be

expected that the expansion of chlorine would be greater than that

of air or of the gases composing it. V. Meyer and Langer (1885)

having remarked that at 1,400° the density of chlorine (taking its

expansion as equal to that of nitrogen) = 29, consider that the

molecules of chlorine split up and partially give molecules Cl, but

it might be maintained that the decrease in density observed only

depends on the increase of the coefficient of expansion.

[9] Investigations on the solubility of chlorine in water (the

solutions evolve all their chlorine on boiling and passing air

through them) show many different peculiarities. First Gay-Lussac,

and subsequently Pelouze, determined that the solubility increases

between 0° and 8°-10° (from 1-1/2 to 2 vols. of chlorine per 100

vols. of water at 0° up to 3 to 2-3/4 at 10°). In the following

note we shall see that this is not due to the breaking-up of the

hydrate at about 8° to 10°, but to its formation below 9°. Roscoe

observed an increase in the solubility of chlorine in the presence

of hydrogen--even in the dark. Berthelot determined an increase of

solubility with the progress of time. Schönbein and others suppose

that chlorine acts on water, forming hypochlorous and hypochloric

acids, (HClO + HCl).

The equilibrium between chlorine and steam as gases and between

water, liquid chlorine, ice, and the solid crystallo-hydrate of

chlorine is evidently very complex. Gibbs, Guldberg (1870) and

others gave a theory for similar states of equilibrium, which was

afterwards developed by Roozeboom (1887), but it would be

inopportune here to enter into its details. It will be sufficient

in the first place to mention that there is now no doubt (according

to the theory of heat, and the direct observations of Ramsay and

Young) that the vapour tensions at one and the same temperature are

different for the liquid and solid states of substances; secondly,

to call attention to the following note; and, thirdly, to state

that, in the presence of the crystallo-hydrate, water between O°·24

and +28°·7 (when the hydrate and a solution may occur

simultaneously) dissolves a different amount of chlorine than it

does in the absence of the crystallo-hydrate.

[10] According to Faraday's data the hydrate of chlorine contains

Cl_{2},10H_{2}O, but Roozeboom (1885) showed that it is poorer in

water and = Cl_{2},8H_{2}O. At first small, almost colourless,

crystals are obtained, but they gradually form (if the temperature

be below their critical point 28°·7, above which they do not

exist) large yellow crystals, like those of potassium chromate.

The specific gravity is 1·23. The hydrate is formed if there be

more chlorine in a solution than it is able to dissolve under the

dissociation pressure corresponding with a given temperature. _In

the presence of the hydrate_ the percentage amount of chlorine at

0° = 0·5, at 9° = 0·9, and at 20° = 1·82. At temperatures below 9°

the solubility (determined by Gay-Lussac and Pelouze, _see_ Note

9) is dependent on the formation of the hydrate; whilst at higher

temperatures under the ordinary pressure the hydrate cannot be

formed, and the solubility of chlorine falls, as it does for all

gases (Chapter I.). If the crystallo-hydrate is not formed, then

below 9° the solubility follows the same rule (6° 1·07 p.c. Cl, 9°

0·95 p.c.). According to Roozeboom, the chlorine evolved by the

hydrate presents the following tensions of dissociation: at 0° =

249 mm., at 4° = 398, at 8° = 620, at 10° = 797, at 14° = 1,400

mm. In this case a portion of the crystallo-hydrate remains solid.

At 9°·6 the tension of dissociation is equal to the atmospheric

pressure. At a higher pressure the crystallo-hydrate may form at

temperatures above 9° up to 28°·7, when the vapour tension of the

hydrate equals the tension of the chlorine. It is evident that the

equilibrium which is established is on the one hand a case of a

complex heterogeneous system, and on the other hand a case of the

solution of solid and gaseous substances in water.

The crystallo-hydrate or chlorine water must be kept in the dark,

or the access of light be prevented by coloured glass, otherwise

oxygen is evolved and hydrochloric acid formed.

Chlorine explodes _with hydrogen_, if a mixture of equal volumes be

exposed to the direct action of the sun's rays[11] or brought into

contact with spongy platinum, or a strongly heated substance, or when

subjected to the action of an electric spark. The explosion in this case

takes place for exactly the same reasons--_i.e._ the evolution of heat

and expansion of the resultant product--as in the case of detonating

gas (Chapter III.) Diffused light acts in the same way, but slowly,

whilst direct sunlight causes an explosion.[12] The hydrochloric acid

gas produced by the reaction of chlorine on hydrogen occupies (at the

original temperature and pressure) a volume equal to the sum of the

original volumes; that is, a reaction of substitution here takes place:

H_{2} + Cl_{2} = HCl + HCl. In this reaction twenty-two thousand heat

units are evolved for one part by weight [1 gram] of hydrogen.[13]

[11] The chemical action of light on a mixture of chlorine and

hydrogen was discovered by Gay-Lussac and Thénard (1809). It has

been investigated by many savants, and especially by Draper,

Bunsen, and Roscoe. Electric or magnesium light, or the light

emitted by the combustion of carbon bisulphide in nitric oxide,

and actinic light in general, acts in the same manner as sunlight,

in proportion to its intensity. At temperatures below -12° light

no longer brings about reaction, or at all events does not give an

explosion. It was long supposed that chlorine that had been

subjected to the action of light was afterwards able to act on

hydrogen in the dark, but it was shown that this only takes place

with moist chlorine, and depends on the formation of oxides of

chlorine. The presence of foreign gases, and even of excess of

chlorine or of hydrogen, very much enfeebles the explosion, and

therefore the experiment is conducted with a detonating mixture

prepared by the action of an electric current on a strong solution

(sp. gr. 1·15) of hydrochloric acid, in which case the water is

not decomposed--that is, no oxygen becomes mixed with the

chlorine.

[12] The quantity of chlorine and hydrogen which combine is

proportional to the intensity of the light--not of all the rays,

but only those so-termed chemical (actinic) rays which produce

chemical action. Hence a mixture of chlorine and hydrogen, when

exposed to the action of light in vessels of known capacity and

surface, may be employed as an actinometer--that is, as a means

for estimating the intensity of the chemical rays, the influence

of the heat rays being previously destroyed, which may be done by

passing the rays through water. Investigations of this kind

(photo-chemical) showed that chemical action is chiefly limited to

the violet end of the spectrum, and that even the invisible

ultra-violet rays produce this action. A colourless gas flame

contains no chemically active rays; the flame coloured green by a

salt of copper evinces more chemical action than the colourless

flame, but the flame brightly coloured yellow by salts of sodium

has no more chemical action than that of the colourless flame.

As the chemical action of light becomes evident in plants,

photography, the bleaching of tissues, and the fading of colours

in the sunlight, and as a means for studying the phenomenon is

given in the reaction of chlorine on hydrogen, this subject has

been the most fully investigated in _photo-chemistry_. The

researches of Bunsen and Roscoe in the fifties and sixties are the

most complete in this respect. Their actinometer contains hydrogen

and chlorine, and is surrounded by a solution of chlorine in

water. The hydrochloric acid is absorbed as it forms, and

therefore the variation in volume indicates the progress of the

combination. As was to be expected, the action of light proved to

be proportional to the time of exposure and intensity of the

light, so that it was possible to conduct detailed photometrical

investigations respecting the time of day and season of the year,

various sources of light, its absorption, &c. This subject is

considered in detail in special works, and we only stop to mention

one circumstance, that a small quantity of a foreign gas decreases

the action of light; for example, 1/330 of hydrogen by 38 p.c.,

1/200 of oxygen by 10 p.c., 1/100 of chlorine by 60 p.c., &c.

According to the researches of Klimenko and Pekatoros (1889), the

photo-chemical alteration of chlorine water is retarded by the

presence of traces of metallic chlorides, and this influence

varies with different metals.

As much heat is evolved in the reaction of chlorine on hydrogen,

and as this reaction, being exothermal, may proceed by itself, the

action of light is essentially the same as that of heat--that is,

it brings the chlorine and hydrogen into the condition necessary

for the reaction--it, as we may say, disturbs the original

equilibrium; this is the work done by the luminous energy. It

seems to me that the action of light on the mixed gases should be

understood in this sense, as Pringsheim (1877) pointed out.

[13] In the formation of steam (from one part by weight [1 gram]

of hydrogen) 29,000 heat units are evolved. The following are the

quantities of heat (thousands of units) evolved in the formation

of various other _corresponding_ compounds of oxygen and of

chlorine (from Thomsen's, and, for Na_{2}O, Beketoff's results):

{2NaCl, 195; CaCl_{2}, 170; HgCl_{2}, 63; 2AgCl, 59.

{ Na_{2}O, 100; CaO, 131; HgO, 42; Ag_{2}O, 6.

{2AsCl_{3}, 143; 2PCl_{5}, 210; CCl_{4}, 21; 2HCl, 44

(gas).

{ As_{2}O_{3}, 155; P_{2}O_{5}, 370; CO_{2}, 97; H_{2}O, 58

(gas).

With the first four elements the formation of the chlorine

compound gives the most heat, and with the four following the

formation of the oxygen compound evolves the greater amount of

heat. The first four chlorides are true salts formed from HCl and

the oxide, whilst the remainder have other properties, as is seen

from the fact that they are not formed from hydrochloric acid and

the oxide, but give hydrochloric acid with water.

These relations show that the affinity of chlorine for hydrogen is

very great and analogous to the affinity between hydrogen and oxygen.

Thus[14] on the one hand by passing a mixture of steam and chlorine

through a red-hot tube, or by exposing water and chlorine to the

sunlight, oxygen is disengaged, whilst on the other hand, as we saw

above, oxygen in many cases displaces chlorine from its compound with

hydrogen, and therefore the reaction H_{2}O + Cl_{2} = 2HCl + O belongs

to the number of reversible reactions, and hydrogen will distribute

itself between oxygen and chlorine. This determines the relation of Cl

to substances containing hydrogen and its reactions in the presence of

water, to which we shall turn our attention after having pointed out the

relation of chlorine to other elements.

[14] This has been already pointed out in Chapter III., Note 5.

Many _metals_ when brought into contact with chlorine immediately

combine with it, and form those metallic chlorides which correspond

with hydrogen chloride and with the oxide of the metal taken. This

combination may proceed rapidly with the evolution of heat and light;

that is, metals are able to burn in chlorine. Thus, for example,

sodium[15] burns in chlorine, synthesising common salt. Metals in

the form of powders burn without the aid of heat, and become highly

incandescent in the process; for instance, antimony, which is a metal

easily converted into a powder.[16] Even such metals as gold and

platinum,[17] which do not combine directly with oxygen and give very

unstable compounds with it, unite directly with chlorine to form

metallic chlorides. Either chlorine water or aqua regia may be employed

for this purpose instead of gaseous chlorine. These dissolve gold and

platinum, converting them into metallic chlorides. _Aqua regia_ is a

mixture of 1 part of nitric acid with 2 to 3 parts of hydrochloric

acid. This mixture converts into soluble chlorides not only those

metals which are acted on by hydrochloric and nitric acids, but also

gold and platinum, which are insoluble in either acid separately.

This action of aqua regia depends on the fact that nitric acid in

acting on hydrochloric acid evolves chlorine. If the chlorine evolved

be transferred to a metal, then a fresh quantity is formed from the

remaining acids and also combines with the metal.[18] Thus the aqua

regia acts by virtue of the chlorine which it contains and disengages.

[15] Sodium remains unaltered in perfectly dry chlorine at the ordinary

temperature, and even when slightly warmed; but the combination is

exceedingly violent at a red heat.

[16] An instructive experiment on combustion in chlorine may be

conducted as follows: leaves of Dutch metal (used instead of gold

for gilding) are placed in a glass globe, and a gas-conducting

tube furnished with a glass cock is placed in the cork closing it,

and the air is pumped out of the globe. The gas-conducting tube is

then connected with a vessel containing chlorine, and the cock

opened; the chlorine rushes in, and the metallic leaves are

consumed.

[17] The behaviour of platinum to chlorine at a high temperature

(1,400°) is very remarkable, because platinous chloride, PtCl_{2},

is then formed, whilst this substance decomposes at a much lower

temperature into chlorine and platinum. Hence, when chlorine comes

into contact with platinum at such high temperatures, it forms

fumes of platinous chloride, and they on cooling decompose, with

the liberation of platinum, so that the phenomenon appears to be

dependent on the volatility of platinum. Deville proved the

formation of platinous chloride by inserting a cold tube inside a

red-hot one (as in the experiment on carbonic oxide). However, V.

Meyer was able to observe the density of chlorine in a platinum

vessel at 1,690°, at which temperature chlorine does not exert

this action on platinum, or at least only to an insignificant

degree.

[18] When left exposed to the air aqua regia disengages chlorine, and

afterwards it no longer acts on gold. Gay-Lussac, in explaining

the action of aqua regia, showed that when heated it evolves,

besides chlorine, the vapours of two chloranhydrides--that of

nitric acid, NO_{2}Cl (nitric acid, NO_{2}OH, in which HO is

replaced by chlorine; _see_ Chapter on Phosphorus), and that of

nitrous acid, NOCl--but these do not act on gold. The formation of

aqua regia may therefore be expressed by 4NHO_{3} + 8HCl =

2NO_{2}Cl + 2NOCl + 6H_{2}O + 2Cl_{2}. The formation of the

chlorides NO_{2}Cl and NOCl is explained by the fact that the

nitric acid is deoxidised, gives the oxides NO and NO_{2}, and

they directly combine with chlorine to form the above anhydrides.

The majority of _non-metals_ also react directly on chlorine; hot

sulphur and phosphorus burn in it and combine with it at the ordinary

temperature. Only nitrogen, carbon, and oxygen do not combine directly

with it. The chlorine compounds formed by the non-metals--for instance,

phosphorus trichloride, PCl_{3}, and sulphurous chloride, &c., do not

have the properties of salts, and, as we shall afterwards see more

fully, correspond to acid anhydrides and acids; for example, PCl_{3}--to

phosphorous acid, P(OH)_{3}:

NaCl FeCl_{2} SnCl_{4} PCl_{3} HCl

Na(HO) Fe(HO)_{2} Sn(HO)_{4} P(HO)_{3} H(HO)

As the above-mentioned relation in composition--_i.e._ substitution of

Cl by the aqueous residue--exists between many chlorine compounds and

their corresponding hydrates, and as furthermore some (acid) hydrates

are obtained from chlorine compounds by the action of water, for

instance,

PCl_{3} + 3H_{2}0 = P(HO)_{3} + 3HCl

Phosphorus Water Phosphorous Hydrochloric

trichloride acid acid

whilst other chlorine compounds are formed from hydroxides and

hydrochloric acid, with the liberation of water, for example,

NaHO + HCl = NaCl + H_{2}O

we endeavour to express this intimate connection between the hydrates

and chlorine compounds by calling the latter _chloranhydrides_. In

general terms, if the hydrate be basic, then,

M(HO) + HCl = MCl + H_{2}O

hydrate + hydrochloric acid = chloranhydride + water

and if the hydrate ROH be acid, then,

RCl + H_{2}O = R(HO) + HCl

Chloranhydride + water = hydrate + hydrochloric acid

The chloranhydrides MCl corresponding to the bases are evidently

metallic chlorides or salts corresponding to HCl. In this manner a

distinct equivalency is marked between the compounds of chlorine and the

so-called hydroxyl radicle (HO), which is also expressed in the analogy

existing between chlorine, Cl_{2}, and hydrogen peroxide, (HO)_{2}.

As regards the chloranhydrides corresponding to acids and non-metals,

they bear but little resemblance to metallic salts. They are nearly all

volatile, and have a powerful suffocating smell which irritates the eyes

and respiratory organs. They react on water like many anhydrides of the

acids, with the evolution of heat and liberation of hydrochloric acid,

forming acid hydrates. For this reason they cannot usually be obtained

from hydrates--that is, acids--by the action of hydrochloric acid, as in

that case water would be formed together with them, and water decomposes

them, converting them into hydrates. There are many intermediate

chlorine compounds between true saline metallic chlorides like sodium

chloride and true acid chloranhydrides, just as there are all kinds of

transitions between bases and acids. Acid chloranhydrides are not only

obtained from chlorine and non-metals, but also from many lower oxides,

by the aid of chlorine. Thus, for example, CO, NO, NO_{2}, SO_{2}, and

other lower oxides which are capable of combining with oxygen may also

combine with a corresponding quantity of chlorine. Thus COCl_{2},

NOCl, NO_{2}Cl, SO_{2}Cl_{2}, &c., are obtained. They correspond with

the hydrates CO(OH)_{2}, NO(OH), NO_{2}(OH), SO_{2}(OH)_{2}, &c., and

to the anhydrides CO_{2}, N_{2}O_{3}, N_{2}O_{5}, SO_{3}, &c. Here

we should notice two aspects of the matter: (1) chlorine combines

with that with which oxygen is able to combine, because it is in many

respects equally if not more energetic than oxygen and replaces it in

the proportion Cl_{2} : O; (2) that highest limit of possible combination

which is proper to a given element or grouping of elements is very

easily and often attained by combination with chlorine. If phosphorus

gives PCl_{3} and PCl_{5}, it is evident that PCl_{5} is the higher

form of combination compared with PCl_{3}. To the form PCl_{5}, or

in general PX_{5}, correspond PH_{4}I, PO(OH)_{3}, POCl_{3}, &c. If

chlorine does not always directly give compounds of the highest possible

forms for a given element, then generally the lower forms combine

with it in order to reach or approach the limit. This is particularly

clear in hydrocarbons, where we see the limit C_{_n_}H_{2_n_+2} very

distinctly. The unsaturated hydrocarbons are sometimes able to combine

with chlorine with the greatest ease and thus reach the limit. Thus

ethylene, C_{2}H_{4}, combines with Cl_{2}, forming the so-called Dutch

liquid or ethylene chloride, C_{2}H_{4}Cl_{2}, because it then reaches

the limit C_{_n_}X_{2_n_+2}. In this and all similar cases the combined

chlorine is able by reactions of substitution to give a hydroxide and

a whole series of other derivatives. Thus a hydroxide called glycol,

C_{2}H_{4}(OH)_{2}, is obtained from C_{2}H_{4}Cl_{2}.

Chlorine _in the presence of water_ very often acts directly _as an

oxidising agent_. A substance A combines with chlorine and gives, for

example, ACl_{2}, and this in turn a hydroxide, A(OH)_{2}, which on

losing water forms AO. Here the chlorine has oxidised the substance

A. This frequently happens in the simultaneous action of water and

chlorine: A + H_{2}O + Cl_{2} = 2HCl + AO. Examples of this oxidising

action of chlorine may frequently be observed both in practical

chemistry and technical processes. Thus, for instance, chlorine in the

presence of water oxidises sulphur and metallic sulphides. In this

case the sulphur is converted into sulphuric acid, and the chlorine

into hydrochloric acid, or a metallic chloride if a metallic sulphide

be taken. A mixture of carbonic oxide and chlorine passed into water

gives carbonic anhydride and hydrochloric acid. Sulphurous anhydride

is oxidised by chlorine in the presence of water into sulphuric acid,

just as it is by the action of nitric acid: SO_{2} + 2H_{2}O + Cl_{2} =

H_{2}SO_{4} + 2HCl.

The oxidising action of chlorine in the presence of water is taken

advantage of in practice for the rapid bleaching of tissues and

fibres. The colouring matter of the fibres is altered by oxidation and

converted into a colourless substance, but the chlorine afterwards acts

on the tissue itself. Bleaching by means of chlorine therefore requires

a certain amount of technical skill in order that the chlorine should

not act on the fibres themselves, but that its action should be limited

to the colouring matter only. The fibre for making writing paper, for

instance, is bleached in this manner. The bleaching property of chlorine

was discovered by Berthollet, and forms an important acquisition to the

arts, because it has in the majority of cases replaced that which before

was the universal method of bleaching--namely, exposure to the sun of

the fabrics damped with water, which is still employed for linens, &c.

Time and great trouble, and therefore money also, have been considerably

saved by this change.[19]

[19] Ozone and peroxide of hydrogen also bleach tissues. As the action

of peroxide of hydrogen is easily controlled by taking a weak

solution, and as it has hardly any action upon the tissues

themselves, it is replacing chlorine more and more as a bleaching

agent. The oxidising property of chlorine is apparent in

destroying the majority of organic tissues, and proves fatal to

organisms. This action of chlorine is taken advantage of in

quarantine stations. But the simple fumigation by chlorine must be

carried on with great care in dwelling places, because chlorine

disengaged into the atmosphere renders it harmful to the health.

The power of chlorine for combination is intimately connected with

its capacity for substitution, because, according to the law of

substitution, if chlorine combines with hydrogen, then it also replaces

hydrogen, and furthermore the combination and substitution are

accomplished in the same quantities. Therefore _the atom of chlorine_

which combines with the atom of hydrogen is also able _to replace the

atom of hydrogen_. We mention this property of chlorine not only because

it illustrates the application of the law of substitution in clear and

historically important examples, but more especially because reactions

of this kind explain those _indirect methods_ of the formation of many

substances which we have often mentioned and to which recourse is had

in many cases in chemistry. Thus chlorine does not act on carbon,[20]

oxygen, or nitrogen, but nevertheless its compounds with these elements

may be obtained by the indirect method of the substitution of hydrogen

by chlorine.

[20] A certain propensity of carbon to attract chlorine is evidenced in

the immense absorption of chlorine by charcoal (Note 7), but, so

far as is at present known (if I am not mistaken, no one has tried

the aid of light), no combination takes place between the chlorine

and carbon.

As chlorine easily combines with hydrogen, and does not act on carbon,

it decomposes hydrocarbons (and many of their derivatives) at a high

temperature, depriving them of their hydrogen and liberating the carbon,

as, for example, is clearly seen when a lighted candle is placed in a

vessel containing chlorine. The flame becomes smaller, but continues

to burn for a certain time, a large amount of soot is obtained, and

hydrochloric acid is formed. In this case the gaseous and incandescent

substances of the flame are decomposed by the chlorine, the hydrogen

combines with it, and the carbon is disengaged as soot.[21] This

action of chlorine on hydrocarbons, &c., proceeds otherwise at lower

temperatures, as we will now consider.

[21] The same reaction takes place under the action of oxygen, with

the difference that it burns the carbon, which chlorine is not

able to do. If chlorine and oxygen compete together at a high

temperature, the oxygen will unite with the carbon, and the

chlorine with the hydrogen.

A very important epoch in the history of chemistry was inaugurated by

the discovery of Dumas and Laurent that chlorine is able to displace

and _replace hydrogen_. This discovery is important from the fact

that chlorine proved to be an element which combines with great ease

simultaneously with both the hydrogen and the element with which the

hydrogen was combined. This clearly proved that there is no opposite

polarity between elements forming stable compounds. Chlorine does not

combine with hydrogen because it has opposite properties, as Dumas and

Laurent stated previously, accounting hydrogen to be electro-positive

and chlorine electro-negative; this is not the reason of their combining

together, for the same chlorine which combines with hydrogen is also

able to replace it without altering many of the properties of the

resultant substance. This substitution of hydrogen by chlorine is termed

_metalepsis_. The mechanism of this substitution is very constant. If we

take a hydrogen compound, preferably a hydrocarbon, and if chlorine acts

directly on it, then there is produced on the one hand hydrochloric acid

and on the other hand a compound containing chlorine in the place of the

hydrogen--so that the chlorine divides itself into two equal portions,

one portion is evolved as hydrochloric acid, and the other portion takes

the place of the hydrogen thus liberated. _Hence this metalepsis is

always accompanied by the formation of hydrochloric acid._[22] The scheme

of the process is as follows:

C_{n}H_{m}X + Cl_{2} = C_{n}H_{m-1}ClX + HCl

Hydrocarbon Free Product of Hydrochloric

chlorine metalepsis acid

Or, in general terms--

RH + Cl_{2} = RCl + HCl.

[22] This division of chlorine into two portions may at the same time

be taken as a clear confirmation of the conception of molecules.

According to Avogadro-Gerhardt's law, the molecule of chlorine (p.

310) contains two atoms of this substance; one atom replaces

hydrogen, and the other combines with it.

The conditions under which metalepsis takes place are also very

constant. In the dark chlorine does not usually act on hydrogen

compounds, but the action commences under the influence of light.

The direct action of the sun's rays is particularly propitious to

metalepsis. It is also remarkable that the presence of traces of certain

substances,[23] especially of iodine, aluminium chloride, antimony

chloride, &c., promotes the action. A trace of iodine added to the

substance subjected to metalepsis often produces the same effect as

sunlight.[24]

[23] Such carriers or media for the transference of chlorine and the

halogens in general were long known to exist in iodine and

antimonious chloride, and have been most fully studied by

Gustavson and Friedel, of the Petroffsky Academy--the former with

respect to aluminium bromide, and the latter with respect to

aluminium chloride. Gustavson showed that if a trace of metallic

aluminium be dissolved in bromine (it floats on bromine, and when

combination takes place much heat and light are evolved), the

latter becomes endowed with the property of entering into

metalepsis, which it is not able to do of its own accord. When

pure, for instance, it acts very slowly on benzene, C_{6}H_{6},

but in the presence of a trace of aluminium bromide the reaction

proceeds violently and easily, so that each drop of the

hydrocarbon gives a mass of hydrobromic acid, and of the product

of metalepsis. Gustavson showed that the _modus operandi_ of this

instructive reaction is based on the property of aluminium bromide

to enter into combination with hydrocarbons and their derivatives.

The details of this and all researches concerning the metalepsis

of the hydrocarbons must be looked for in works on organic

chemistry.

[24] As small admixtures of iodine, aluminium bromide, &c., aid the

metalepsis of large quantities of a substance, just as nitric

oxide aids the reaction of sulphurous anhydride on oxygen and

water, so the principle is essentially the same in both cases.

Effects of this kind (which should also be explained by a chemical

reaction proceeding at the surfaces) only differ from true contact

phenomena in that the latter are produced by solid bodies and are

accomplished at their surfaces, whilst in the former all is in

solution. Probably the action of iodine is founded on the

formation of iodine chloride, which reacts more easily than

chlorine.

If marsh gas be mixed with chlorine and the mixture ignited, then the

hydrogen is entirely taken up from the marsh gas and hydrochloric

acid and carbon formed, but there is no metalepsis.[25] But if a

mixture of equal volumes of chlorine and marsh gas be exposed to the

action of diffused light, then the greenish yellow mixture gradually

becomes colourless, and hydrochloric acid and the first product of

metalepsis--namely, methyl chloride--are formed:

CH_{4} + Cl_{2} = CH_{3}Cl + HCl

Marsh gas Chlorine Methyl chloride Hydrochloric acid

[25] Metalepsis belongs to the number of delicate reactions--if it may

be so expressed--as compared with the energetic reaction of

combustion. Many cases of substitution are of this kind. Reactions

of metalepsis are accompanied by an evolution of heat, but in a

less quantity than that evolved in the formation of the resulting

quantity of the halogen acids. Thus the reaction C_{2}H_{6} +

Cl_{2} = C_{2}H_{5}Cl + HCl, according to the data given by

Thomsen, evolves about 20,000 heat units, whilst the formation of

hydrochloric acid evolves 22,000 units.

The volume of the mixture remains unaltered. The methyl chloride which

is formed is a gas. If it be separated from the hydrochloric acid (it

is soluble in acetic acid, in which hydrochloric acid is but sparingly

soluble) and be again mixed with chlorine, then it may be subjected

to a further metalepsical substitution--the second atom of hydrogen

may be substituted by chlorine, and a liquid substance, CH_{2}Cl_{2},

called methylene chloride, will be obtained. In the same manner

the substitution may be carried on still further, and CHCl_{3}, or

chloroform, and lastly carbon tetrachloride, CCl_{4}, will be produced.

Of these substances the best known is chloroform, owing to its being

formed from many organic substances (by the action of bleaching powder)

and to its being used in medicine as an anæsthetic; chloroform boils

at 62° and carbon tetrachloride at 78°. They are both colourless

odoriferous liquids, heavier than water. The progressive substitution of

hydrogen by chlorine is thus evident, and it can be clearly seen that

the double decompositions are accomplished between molecular quantities

of the substance--that is, between equal volumes in a gaseous state.

_Carbon tetrachloride_, which is obtained by the metalepsis of marsh

gas, cannot be obtained directly from chlorine and carbon, but it

may be obtained from certain compounds of carbon--for instance, from

carbon bisulphide--if its vapour mixed with chlorine be passed through

a red-hot tube. Both the sulphur and carbon then combine with the

chlorine. It is evident that by ultimate metalepsis a corresponding

carbon chloride may be obtained from any hydrocarbon--indeed, the number

of chlorides of carbon C_{_n_}Cl_{2_m_} already known is very large.

As a rule, the fundamental chemical characters of hydrocarbons are not

changed by metalepsis; that is, if a neutral substance be taken, then

the product of metalepsis is also a neutral substance, or if an acid

be taken the product of metalepsis also has acid properties. Even the

crystalline form not unfrequently remains unaltered after metalepsis.

The metalepsis of acetic acid, CH_{3}·COOH, is historically the most

important. It contains three of the atoms of the hydrogen of marsh gas,

the fourth being replaced by carboxyl, and therefore by the action of

chlorine it gives three products of metalepsis (according to the amount

of the chlorine and conditions under which the reaction takes place),

mono-, di-, and tri-chloracetic acids--CH_{2}Cl·COOH, CHCl_{2}·COOH, and

CCl_{3}·COOH; they are all, like acetic acid, monobasic. The resulting

products of metalepsis, in containing an element which so easily acts

on metals as chlorine, possess the possibility of attaining a further

complexity of molecules of which the original hydrocarbon is often in

no way capable. Thus on treating with an alkali (or first with a salt

and then with an alkali, or with a basic oxide and water, &c.) the

chlorine forms a salt with its metal, and the hydroxyl radicle takes

the place of the chlorine--for example, CH_{3}·OH is obtained from

CH_{3}Cl. By the action of metallic derivatives of hydrocarbons--for

example, CH_{3}Na--the chlorine also gives a salt, and the hydrocarbon

radicle--for instance, CH_{3}--takes the place of the chlorine. In

this, or in a similar manner, CH_{3}·CH_{3}, or C_{2}H_{6} is obtained

from CH_{3}Cl and C_{6}H_{5}·CH_{3} from C_{6}H_{6}. The products of

metalepsis also often react on ammonia, forming hydrochloric acid (and

thence NH_{4}Cl) and an amide; that is, the product of metalepsis,

with the ammonia radicle NH_{2}, &c. in the place of chlorine. Thus

by means of metalepsical substitution methods were found in chemistry

for an artificial and general means of the formation of complex carbon

compounds from more simple compounds which are often totally incapable

of direct reaction. Besides which, this key opened the doors of that

secret edifice of complex organic compounds into which man had up to

then feared to enter, supposing the hydrocarbon elements to be united

only under the influence of those mystic forces acting in organisms.[26]

[26] With the predominance of the representation of compound radicles

(this doctrine dates from Lavoisier and Gay-Lussac) in organic

chemistry, it was a very important moment in its history when it

became possible to gain an insight into the structure of the

radicles themselves. It was clear, for instance, that ethyl,

C_{2}H_{5}, or the radicle of common alcohol, C_{2}H_{5}·OH,

passes, without changing, into a number of ethyl derivatives, but

its relation to the still simpler hydrocarbons was not clear, and

occupied the attention of science in the 'forties' and 'fifties.'

Having obtained ethyl hydride, C_{2}H_{5}H = C_{2}H_{6}, it was

looked on as containing the same ethyl, just as methyl hydride,

CH_{4} = CH_{3}H, was considered as existing in methane. Having

obtained free methyl, CH_{3}CH_{3} = C_{2}H_{6}, from it, it was

considered as a derivative of methyl alcohol, CH_{3}OH, and as

only isomeric with ethyl hydride. By means of the products of

metalepsis it was proved that this is not a case of isomerism but

of strict identity, and it therefore became clear that ethyl is

methylated methyl, C_{2}H_{5} = CH_{2}CH_{3}. In its time a still

greater impetus was given by the study of the reactions of

monochloracetic acid, CH_{2}Cl·COOH, or CO(CH_{2}Cl)(OH). It

appeared that metalepsical chlorine, like the chlorine of

chloranhydrides--for instance, of methyl chloride, CH_{3}Cl, or

ethyl chloride, C_{2}H_{5}Cl--is capable of substitution; for

example, glycollic acid, CH_{2}(OH)(CO_{2}H), or

CO(CH_{2}·OH)(OH), was obtained from it, and it appeared that the

OH in the group CH_{2}(OH) reacted like that in alcohols, and it

became clear, therefore, that it was necessary to examine the

radicles themselves by analysing them from the point of view of

the bonds connecting the constituent atoms. Whence arose the

present doctrine of the structure of the carbon compounds. (_See_

Chapter VIII., Note 42.)

It is not only hydrocarbons which are subject to metalepsis. Certain

other hydrogen compounds, under the action of chlorine, also give

corresponding chlorine derivatives in exactly the same manner; for

instance, ammonia, caustic potash, caustic lime, and a whole series of

_alkaline_ substances.[27] In fact, just as the hydrogen in marsh gas

can be replaced by chlorine and form methyl chloride, so the hydrogen

in caustic potash, KHO, ammonia, NH_{3}, and calcium hydroxide,

CaH_{2}O_{2} or Ca(OH)_{2}, may be replaced by chlorine and give

potassium hypochlorite, KClO, calcium hypochlorite, CaCl_{2}O_{2},

and the so-called chloride of nitrogen, NCl_{3}. For not only is the

correlation in composition the same as in the substitution in marsh gas,

but the whole mechanism of the reaction is the same. Here also two atoms

of chlorine act: one takes the place of the hydrogen whilst the other is

evolved as hydrochloric acid, only in the former case the hydrochloric

acid evolved remained free, and in the latter, in presence of alkaline

substances, it reacts on them. Thus, in the action of chlorine on

caustic potash, the hydrochloric acid formed acts on another quantity of

caustic potash and gives potassium chloride and water, and therefore not

only KHO + Cl_{2} = HCl + KClO, but also KHO + HCl = H_{2}O + KCl, and

the result of both simultaneous phases will be 2KHO + Cl_{2} = H_{2}O +

KCl + KClO. We will here discuss certain special cases.

[27] By including many instances of the action of chlorine under

metalepsis we not only explain the indirect formation of CCl_{4},

NCl_{3}, and Cl_{2}O by one method, but we also arrive at the fact

that the reactions of the metalepsis of the hydrocarbons lose that

exclusiveness which was often ascribed to them. Also by subjecting

the chemical representations to the law of substitution we may

foretell metalepsis as a particular case of a general law.

The action of chlorine on ammonia may either result in the entire

breaking up of the ammonia, with the evolution of gaseous nitrogen, or

in a product of metalepsis (as with CH_{4}). With an excess of chlorine

and the aid of heat the ammonia is decomposed, with the disengagement

of free nitrogen.[28] This reaction evidently results in the formation

of sal-ammoniac, 8NH_{3} + 3Cl_{2} = 6NH_{4}Cl + N_{2}. But if the

ammonium salt be in excess, then the reaction takes the direction of the

replacement of the hydrogen in the ammonia by chlorine. The principal

result is that NH_{3} + 3Cl_{2} forms NCl_{3} + 3HCl.[29] The resulting

product of metalepsis, or _chloride of nitrogen_, NCl_{3}, discovered by

Dulong, is a liquid having the property of decomposing with excessive

ease not only when heated, but even under the action of mechanical

influences, as by a blow or by contact with certain solid substances.

The explosion which accompanies the decomposition is due to the fact

that the liquid chloride of nitrogen gives gaseous products, nitrogen

and chlorine.[29 bis]

[28] This may be taken advantage of in the preparation of nitrogen.

If a large excess of chlorine water be poured into a beaker, and a

small quantity of a solution of ammonia be added, then, after

shaking, nitrogen is evolved. If chlorine act on a dilute solution

of ammonia, the volume of nitrogen does not correspond with the

volume of the chlorine taken, because ammonium hypochlorite is

formed. If ammonia gas be passed through a fine orifice into a

vessel containing chlorine, the reaction of the formation of

nitrogen is accompanied by the emission of light and the

appearance of a cloud of sal-ammoniac. In all these instances an

excess of chlorine must be present.

[29] The hydrochloric acid formed combines with ammonia, and therefore

the final result is 4NH_{3} + 3Cl_{2} = NCl_{3} + 3NH_{4}Cl. For

this reason, more ammonia must enter into the reaction, but the

metalepsical reaction in reality only takes place with an excess

of ammonia or its salt. If bubbles of chlorine be passed through a

fine tube into a vessel containing ammonia gas, each bubble gives

rise to an explosion. If, however, chlorine be passed into a

solution of ammonia, the reaction at first brings about the

formation of nitrogen, because chloride of nitrogen acts on

ammonia like chlorine. But when sal-ammoniac has begun to form,

then the reaction directs itself towards the formation of chloride

of nitrogen. The first action of chlorine on a solution of

sal-ammoniac always causes the formation of chloride of nitrogen,

which then reacts on ammonia thus: NCl_{3} + 4NH_{3} = N_{2} +

3NH_{4}Cl. Therefore, so long as the liquid is alkaline from the

presence of ammonia the chief product will be nitrogen. The

reaction NH_{4}Cl + 3Cl_{2} = NCl_{3} + 4HCl is reversible; with a

dilute solution it proceeds in the above-described direction

(perhaps owing to the affinity of the hydrochloric acid for the

excess of water), but with a strong solution of hydrochloric acid

it takes the opposite direction (probably by virtue of the

affinity of hydrochloric acid for ammonia). Therefore there must

exist a very interesting case of equilibrium between ammonia,

hydrochloric acid, chlorine, water, and chloride of nitrogen which

has not yet been investigated. The reaction NCl_{3} + 4HCl =

NH_{4}Cl + 3Cl_{2} enabled Deville and Hautefeuille to determine

the composition of chloride of nitrogen. When slowly decomposed by

water, chloride of nitrogen gives, like a chloranhydride, nitrous

acid or its anhydride, 2NCl_{3} + 3H_{2}O = N_{2}O_{3} + 6HCl.

From these observations it is evident that chloride of nitrogen

presents great chemical interest, which is strengthened by its

analogy with trichloride of phosphorus. The researches of F. F.

Selivanoff (1891-94) prove that NCl_{3} may be regarded as an

ammonium derivative of hypochlorous acid. Chloride of nitrogen is

decomposed by dilute sulphuric acid in the following manner:

NCl_{3} + 3H_{2}O + H_{2}SO_{4} = NH_{4}HSO_{4} + 3HClO. This

reaction is reversible and is only complete when some substance,

combining with HClO (for instance, succinimide) or decomposing it,

is added to the liquid. This is easily understood from the fact

that hypochlorous acid itself, HClO, may, according to the view

held in this book, be regarded as the product of the metalepsis of

water, and consequently bears the same relation to NCl_{3} as

H_{2}O does to NH_{3}, or as RHO to RNH_{2}, R_{2}NH, and

R_{3}N--that is to say, NCl_{3} corresponds as an ammonium

derivative to ClOH and Cl_{2} in exactly the same manner as NR_{3}

corresponds to ROH and R_{2}. The connection of NCl_{3} and other

similar explosive chloro-nitrogen compounds (called chloryl

compounds by Selivanoff; for example, the C_{2}H_{5}NCl_{2} of

Wurtz is chloryl ethylamine), such as NRCl_{2} (as

NC_{2}H_{5}Cl_{2}), and NR_{2}Cl (for instance, N(CH_{3}CO)HCl,

chlorylacetamide, and N(C_{2}H_{5})_{2}Cl, chloryl diethylamine)

with HClO is evident from the fact that under certain

circumstances these compounds give hypochlorous acid, with water,

for instance, NR_{2}Cl + H_{2}O = NR_{2}H + HClO, and frequently

act (like NCl_{3} and HClO, or Cl_{2}) in an oxidising and

chloridising manner. We may take chloryl succinimide,

C_{2}H_{4}(CO)_{2}NCl for example. It was obtained by Bender by

the action of HClO upon succinimide, C_{2}H_{4}(CO)_{2}NH, and is

decomposed by water with the re-formation of amide and HClO (the

reaction is reversible). Selivanoff obtained, investigated, and

classified many of the compounds NR_{2}Cl and NRCl_{2}, where R is

a residue of organic acids or alcohols, and showed their

distinction from the chloranhydrides, and thus supplemented the

history of chloride of nitrogen, which is the simplest of the

amides containing chlorine, NR_{3}, where R is fully substituted

by chlorine.

[29 bis] In preparing NCl_{3} every precaution must he used to guard

against an explosion, and care should he taken that the NCl_{3}

remains under a layer of water. Whenever an ammoniacal substance

comes into contact with chlorine great care must be taken, because

it may be a case of the formation of such products and a very

dangerous explosion may ensue. The liquid product of the

metalepsis of ammonia may be most safely prepared in the form of

small drops by the action of a galvanic current on a slightly warm

solution of sal-ammoniac; chlorine is then evolved at the positive

pole, and this chlorine acting on the ammonia gradually forms the

product of metalepsis which floats on the surface of the liquid

(being carried up by the gas), and if a layer of turpentine be

poured on to it these small drops, on coming into contact with the

turpentine, give feeble explosions, which are in no way dangerous

owing to the small mass of the substance formed. Drops of chloride

of nitrogen may with great caution be collected for investigation

in the following manner. The neck of a funnel is immersed in a

basin containing mercury, and first a saturated solution of common

salt is poured into the funnel, and above it a solution of

sal-ammoniac in 9 parts of water. Chlorine is then slowly passed

through the solutions, when drops of chloride of nitrogen fall

into the salt water.

Chloride of nitrogen is a yellow oily liquid of sp. gr. 1·65, which

boils at 71°, and breaks up into N + Cl_{3} at 97°. The contact of

phosphorus, turpentine, india-rubber, &c. causes an explosion, which

is sometimes so violent that a small drop will pierce through a thick

board. The great ease with which chloride of nitrogen decomposes is

dependent upon the fact that it is formed with an absorption of heat,

which it evolves when decomposed, to the amount of about 38,000 heat

units for NCl_{3}, as Deville and Hautefeuille determined.

Chlorine, when absorbed by a solution of caustic soda (and also of other

alkalis) at the ordinary temperature, causes the replacement of the

hydrogen in the caustic soda by the chlorine, with the formation of

sodium chloride by the hydrochloric acid, so that the reaction may be

represented in two phases, as described above. In this manner, sodium

hypochlorite, NaClO, and sodium chloride are simultaneously formed: 2NaHO

+ Cl_{2} = NaCl + NaClO + H_{2}O. The resultant solution contains NaClO

and is termed 'eau de Javelle.' An exactly similar reaction takes place

when chlorine is passed over dry hydrate of lime at the ordinary

temperature: 2Ca(HO)_{2} + 2Cl_{2} = CaCl_{2}O_{2} + CaCl_{2} + 2H_{2}O.

A mixture of the product of metalepsis with calcium chloride is obtained.

This mixture is employed in practice on a large scale, and is termed

'bleaching powder,' owing to its acting, especially when mixed with

acids, as a bleaching agent on tissues, so that it resembles chlorine in

this respect. It is however preferable to chlorine, because the

destructive action of the chlorine can be moderated in this case, and

because it is much more convenient to deal with a solid substance than

with gaseous chlorine. Bleaching powder is also called _chloride of

lime_, because it is obtained from chlorine and hydrate of lime, and

contains[30] both these substances. It may be prepared in the laboratory

by passing a current of chlorine through a cold mixture of water and lime

(milk of lime). The mixture must be kept cold, as otherwise 3Ca(ClO)_{2}

passes into 2CaCl_{2} + Ca(ClO_{3})_{2}. In the manufacture of bleaching

powder in large quantities at chemical works, the purest possible slaked

lime is taken and laid in a thin layer in large flat chambers, M (whose

walls are made of Yorkshire flags or tarred wood, on which chlorine has

no action), and into which chlorine gas is introduced by lead tubes. The

distribution of the plant is shown in the annexed drawing (fig. 67).

[Illustration: FIG. 67.--Apparatus for the manufacture of bleaching

powder (on a small scale) by the action of chlorine, which is generated

in the vessels C, on lime, which is charged into M.]