Chapter I: , because in many instances unstable highly iodised (1)

compounds, resembling crystallo-hydrates, have been obtained from
such solutions. Thus iodide of tetramethylammonium,
N(CH_{3})_{4}I, combines with I_{2}, and I_{4}. Even a solution of
iodine in a saturated solution of potassium iodide presents
indications of the formation of a definite compound KI_{3}. Thus,
an alcoholic solution of KI_{3} does not give up iodine to carbon
bisulphide, although this solvent takes up iodine from an
alcoholic solution of iodine itself (Girault, Jörgensen, and
others). The instability of these compounds resembles the
instability of many crystallo-hydrates, for instance of
HCl,2H_{2}O.

If we compare the four elements, fluorine, chlorine, bromine, and iodine, we see in them an example of analogous substances which arrange themselves by their physical properties in the same order as they stand in respect to their atomic and molecular weights. If the weight of the molecule be large, the substance has a higher specific gravity, a higher melting and boiling point, and a whole series of properties depending on this difference in its fundamental properties. Chlorine in a free state boils at about -35°, bromine boils at 60°, and iodine only above 180°. According to Avogadro-Gerhardt's law, the vapour densities of these elements in a gaseous state are proportional to their atomic weights, and here, at all events approximately, the densities in a liquid (or solid) state are also almost in the ratio of their atomic weights. Dividing the atomic weight of chlorine (35·5) by its specific gravity in a liquid state (1·3), we obtain a volume = 27, for bromine (80/3·1) 26, and for iodine also (127/4·9) 26.[64]

[64] The equality of the atomic volumes of the halogens themselves is
all the more remarkable because in all the halogen compounds the
volume augments with the substitution of fluorine by chlorine,
bromine, and iodine. Thus, for example, the volume of sodium
fluoride (obtained by dividing the weight expressed by its formula
by its specific gravity) is about 15, of sodium chloride 27, of
sodium bromide 32, and of sodium iodide 41. The volume of silicon
chloroform, SiHCl_{3}, is 82, and those of the corresponding
bromine and iodine compounds are 108 and 122 respectively. The
same difference also exists in solutions; for example, NaCl +
200H_{2}O has a sp. gr. (at 15°/4°) of 1·0106, consequently the
volume of the solution 3,658·5/1·0106 = 3,620, hence the volume of
sodium chloride in solution = 3,620-3,603 (this is the volume of
200 H_{2}O) = 17, and in similar solutions, NaBr = 26 and NaI =
35.

The metallic bromides and iodides are in the majority of cases, in most respects analogous to the corresponding chlorides,[65] but chlorine displaces the bromine and iodine from them, and bromine liberates iodine from iodides, which is taken advantage of in the preparation of these halogens. However, the researches of Potilitzin showed that a _reverse_ displacement of chlorine by bromine may occur both in solutions and in ignited metallic chlorides in an atmosphere of bromine vapour--that is, a distribution of the metal (according to Berthollet's doctrine) takes place between the halogens, although however the larger portion, still unites with the chlorine, which shows its greater affinity for metals as compared with that of bromine and iodine.[66] The latter, however, sometimes behave with respect to metallic oxides in exactly the same manner as chlorine. Gay-Lussac, by igniting potassium carbonate in iodine vapour, obtained (as with chlorine) an evolution of oxygen and carbonic anhydride, K_{2}CO_{3} + I_{2} = 2KI + CO_{2} + O, only the reactions between the halogens and oxygen are more easily reversible with bromine and iodine than with chlorine. Thus, at a red heat oxygen displaces iodine from barium iodide. Aluminium iodide burns in a current of oxygen (Deville and Troost), and a similar, although not so clearly marked, relation exists for aluminium chloride, and shows that the halogens have a distinctly smaller affinity for those metals which only form feeble bases. This is still more the case with the non-metals, which form acids and evolve much more heat with oxygen than with the halogens (Note 13). But in all these instances the affinity (and amount of heat evolved) of iodine and bromine is less than that of chlorine, probably because the atomic weights are greater. The smaller store of energy in iodine and bromine is seen still more clearly in the relation of the halogens to hydrogen. In a gaseous state they all enter, with more or less ease, into direct combination with gaseous hydrogen--for example, in the presence of spongy platinum, forming halogen acids, HX--but the latter are far from being equally stable; hydrogen chloride is the most stable, hydrogen iodide the least so, and hydrogen bromide occupies an intermediate position. A very high temperature is required to decompose hydrogen chloride even partially, whilst hydrogen iodide is decomposed by light even at the ordinary temperature and very easily by a red heat. Hence the reaction I_{2} + H_{2} = HI + HI is very easily reversible, and consequently has a limit, and hydrogen iodide easily dissociates.[67] Judging by the direct measurement of the heat evolved (22,000 heat units) in the formation of HCl, the conversion of 2HCl into H_{2} + Cl_{2} requires the expenditure of 44,000 heat units. The decomposition of 2HBr into H_{2} + Br_{2} only requires, if the bromine be obtained in a gaseous state, a consumption of about 24,000 units, whilst in the decomposition of 2HI into H_{2} + I_{2} as vapour about 3,000 heat units are _evolved_;[68] these facts, without doubt, stand in causal connection with the great stability of hydrogen chloride, the easy decomposability of hydrogen iodide, and the intermediate properties of hydrogen bromide. From this it would be expected that chlorine is capable of decomposing water with the evolution of oxygen, whilst iodine has not the energy to produce this disengagement,[69] although it is able to liberate the oxygen from the oxides of potassium and sodium, the affinity of these metals for the halogens being very considerable. For this reason oxygen, especially in compounds from which it can be evolved readily (for instance, ClHO, CrO_{3}, &c.), easily decomposes hydrogen iodide. A mixture of hydrogen iodide and oxygen burns in the presence of an ignited substance, forming water and iodine. Drops of nitric acid in an atmosphere of hydrogen iodide cause the disengagement of violet fumes of iodine and brown fumes of nitric peroxide. In the presence of alkalis and an excess of water, however, iodine is able to effect oxidation like chlorine--that is, it decomposes water; the action is here aided by the affinity of hydrogen iodide for the alkali and water, just as sulphuric acid helps zinc to decompose water. But the relative instability of hydriodic acid is best seen in comparing the acids in a gaseous state. If the halogen acids be dissolved in water, they evolve so much heat that they approach much nearer to each other in properties. This is seen from thermochemical data, for in the formation of HX in solution (in a large excess of water) from the _gaseous_ elements there is _evolved_ for HCl 39,000, for HBr 32,000, and for HI 18,000 heat units.[70] But it is especially evident from the fact that solutions of hydrogen bromide and iodide in water have many points in common with solutions of hydrogen chloride, both in their capacity to form hydrates and fuming solutions of constant boiling point, and in their capacity to form haloid salts, &c. by reacting on bases.

[65] But the density (and also molecular volume, Note 64) of a bromine
compound is always greater than that of a chlorine compound,
whilst that of an iodine compound is still greater. The order is
the same in many other respects. For example, an iodine compound
has a higher boiling point than a bromine compound, &c.

[66] A. L. Potilitzin showed that in heating various metallic chlorides
in a closed tube, with an equivalent quantity of bromine, a
distribution of the metal between the halogens always occurs, and
that the amounts of chlorine replaced by the bromine in the
ultimate product are proportional to the atomic weights of the
metals taken and inversely proportional to their equivalence.
Thus, if NaCl + Br be taken, then out of 100 parts of chlorine,
5·54 are replaced by the bromine, whilst with AgCl + Br 27·28
parts are replaced. These figures are in the ratio 1 : 4·9, and
the atomic weights Na : Ag = 1 : 4·7. In general terms, if a
chloride MCl_{_n_} be taken, it gives with _n_Br a percentage
substitution = 4M/_n_^2 where M is the atomic weight of the metal.
This law was deduced from observations on the chlorides of Li, K,
Na, Ag (_n_ = 1), Ca, Sr, Ba, Co, Ni, Hg, Pb (_n_ = 2), Bi (_n_ =
3), Sn (_n_ = 4), and Fe_{2} (_n_ = 6).

In these determinations of Potilitzin we see not only a brilliant
confirmation of Berthollet's doctrine, but also the first effort
to directly determine the affinities of elements by means of
displacement. The chief object of these researches consisted in
proving whether a displacement occurs in those cases where heat is
absorbed, and in this instance it should be absorbed, because the
formation of all metallic bromides is attended with the evolution
of less heat than that of the chlorides, as is seen by the figures
given in Note 55.

If the mass of the bromine be increased, then the amount of
chlorine displaced also increases. For example, if masses of
bromine of 1 and 4 equivalents act on a molecule of sodium
chloride, then the percentages of the chlorine displaced will be
6·08 p.c. and 12·46 p.c.; in the action of 1, 4, 25, and 100
molecules of bromine on a molecule of barium chloride, there will
be displaced 7·8, 17·6, 35·0, and 45·0 p.c. of chlorine. If an
equivalent quantity of hydrochloric acid act on metallic bromides
in closed tubes, and in the absence of water at a temperature of
300°, then the percentages of the substitution of the bromine by
the chlorine in the double decomposition taking place between
univalent metals are inversely proportional to their atomic
weights. For example, NaBr + HCl gives at the limit 21 p.c. of
displacement, KCl 12 p.c. and AgCl 4-1/4 p.c. Essentially the same
action takes place in an aqueous solution, although the phenomenon
is complicated by the participation of the water. The reactions
proceed spontaneously in one or the other direction at the
ordinary temperature but at different _rates_. In the action of a
dilute solution (1 equivalent per 5 litres) of sodium chloride on
silver bromide at the ordinary temperature the amount of bromine
replaced in six and a half days is 2·07 p.c., and with potassium
chloride 1·5 p.c. With an excess of the chloride the magnitude of
the substitution increases. These conversions also proceed with
the absorption of heat. The reverse reactions evolving heat
proceed incomparably more rapidly, but also to a certain limit;
for example, in the reaction AgCl + RBr the following percentages
of silver bromide are formed in different times:

hours 2 3 22 96 120
K 79·82 87·4 88·22 -- 94·21
Na 83·63 90·74 91·70 95·49 --

That is, the conversions which are accompanied by an evolution of
heat proceed with very much greater rapidity than the reverse
conversions.

[67] _The dissociation of hydriodic acid_ has been studied in detail by
Hautefeuille and Lemoine, from whose researches we extract the
following information. The decomposition of hydriodic acid is
decided, but proceeds slowly at 180°; the rate and limit of
decomposition increase with a rise of temperature. The reverse
action--that is, I_{2} + H_{2} = 2HI--proceeds not only under the
influence of spongy platinum (Corenwinder), which also accelerates
the decomposition of hydriodic acid, but also by itself, although
slowly. The limit of the reverse reaction remains the same with or
without spongy platinum. An increase of pressure has a very
powerful accelerative effect on the rate of formation of hydriodic
acid, and therefore spongy platinum by condensing gases has the
same effect as increase of pressure. At the atmospheric pressure
the decomposition of hydriodic acid reaches the limit at 250° in
several months, and at 440° in several hours. The limit at 250° is
about 18 p.c. of decomposition--that is, out of 100 parts of
hydrogen previously combined in hydriodic acid, about 18 p.c. may
be disengaged at this temperature (this hydrogen may be easily
measured, and the measure of dissociation determined), but not
more; the limit at 440° is about 26 p.c. If the pressure under
which 2HI passes into H_{2} + I_{2} be 4-1/2 atmospheres, then the
limit is 24 p.c.; under a pressure of 1/5 atmosphere the limit is
29 p.c. The small influence of pressure on the dissociation of
hydriodic acid (compared with N_{2}O_{4}, Chapter VI. Note 46) is
due to the fact that the reaction 2HI = I_{2} + H_{2} is not
accompanied by a change of volume. In order to show the influence
of time, we will cite the following figures referring to 350°: (1)
Reaction H_{2} + I_{2}; after 3 hours, 88 p.c. of hydrogen remained
free; 8 hours, 69 p.c.; 34 hours, 48 p.c.; 76 hours, 29 p.c.; and
327 hours, 18·5 p.c. (2) The reverse decomposition of 2HI; after 9
hours, 3 p.c. of hydrogen was set free, and after 250 hours 18·6
p.c.--that is, the limit was reached. The addition of extraneous
hydrogen diminishes the limit of the reaction of decomposition, or
increases the formation of hydriodic acid from iodine and
hydrogen, as would be expected from Berthollet's doctrine (Chapter
X.). Thus at 440° 26 p.c. of hydriodic acid is decomposed if there
be no admixture of hydrogen, while if H_{2} be added, then at the
limit only half as large a mass of HI is decomposed. Therefore, if
an infinite mass of hydrogen be added there will be no
decomposition of the hydriodic acid. Light aids the decomposition
of hydriodic acid very powerfully. At the ordinary temperature 80
p.c. is decomposed under the influence of light, whilst under the
influence of heat alone this limit corresponds with a very high
temperature. The distinct action of light, spongy platinum, and of
impurities in glass (especially of sodium sulphate, which
decomposes hydriodic acid), not only render the investigations
difficult, but also show that in reactions like 2HI = I_{2} + H_{2},
which are accompanied by slight heat effects, all foreign and
feeble influences may strongly affect the progress of the action
(Note 47).

[68] The thermal determinations of Thomsen (at 18°) gave in thousands
of calories, Cl + H = +22, HCl + Aq (that is, on dissolving HCl in
a large amount of water) = +17·3, and therefore H + Cl + Aq =
+39·3. In taking molecules, all these figures must be doubled. Br
+ H = +8·4; HBr + Aq = 19·9; H + Br + Aq = +28·3. According to
Berthelot 7·2 are required for the vaporisation of Br_{2}, hence
Br_{2} + H_{2} = 16·8 + 7·2 = +24, if Br_{2} be taken as vapour
for comparison with Cl_{2}. H + I =-6·0, HI + Aq = 19·2; H + I +
Aq= +13·2, and, according to Berthelot, the heat of fusion of
I_{2} = 3·0, and of vaporisation 6·0 thousand heat units, and
therefore I_{2} + H_{2} =-2(6·0) + 3 + 6 =-3·0, if the iodine be
taken as vapour. Berthelot, on the basis of his determinations,
gives, however, +0·8 thousand heat units. Similar contradictory
results are often met with in thermochemistry owing to the
imperfection of the existing methods, and particularly the
necessity of depending on indirect methods for obtaining the
fundamental figures. Thus Thomsen decomposed a dilute solution of
potassium iodide by gaseous chlorine; the reaction gave +26·2,
whence, having first determined the heat effects of the reactions
KHO + HCl, KHO + HI and Cl + H in aqueous solutions, it was
possible to find H + I + Aq; then, knowing HI + Aq, to find I + H.
It is evident that unavoidable errors may accumulate.

[69] One can believe, however, on the basis of Berthollet's doctrine,
and the observations of Potilitzin (Note 66), that a certain slow
decomposition of water by iodine takes place. On this view the
observations of Dossios and Weith on the fact that the solubility
of iodine in water increases after the lapse of several months
will be comprehensible. Hydriodic acid is then formed, and it
increases the solubility. If the iodine be extracted from such a
solution by carbon bisulphide, then, as the authors showed, after
the action of nitrous anhydride iodine may be again detected in
the solution by means of starch. It can easily be understood that
a number of similar reactions, requiring much time and taking
place in small quantities, have up to now eluded the attention of
investigators, who even still doubt the universal application of
Berthollet's doctrine, or only see the thermochemical side of
reactions, or else neglect to pay attention to the element of time
and the influence of mass.

[70] On the basis of the data in Note 68.

In consequence of what has been said above, it follows that _hydrobromic and hydriodic acids_, being substances which are but slightly stable, cannot be evolved in a gaseous state under many of those conditions under which hydrochloric acid is formed. Thus if sulphuric acid in solution acts on sodium iodide, all the same phenomena take place as with sodium chloride (a portion of the sodium iodide gives hydriodic acid, and all remains in solution), but if sodium iodide be mixed with strong sulphuric acid, then the oxygen of the latter decomposes the hydriodic acid set free, with liberation of iodine, H_{2}SO_{4} + 2HI = 2H_{2}O + SO_{2} + I_{2}. This reaction takes place in the reverse direction in the presence of a _large quantity_ of water (2,000 parts of water per 1 part of SO_{2}), in which case not only the affinity of hydriodic acid for water is brought to light but also the action of water in directing chemical reactions in which it participates.[71] Therefore, with a halogen salt, it is easy to obtain gaseous hydrochloric acid by the action of sulphuric acid, but neither hydrobromic nor hydriodic acid can be so obtained in the free state (as gases).[72] Other methods have to be resorted to for their preparation, and recourse must not be had to compounds of oxygen, which are so easily able to destroy these acids. Therefore hydrogen sulphide, phosphorus, &c., which themselves easily take up oxygen, are introduced as means for the conversion of bromine and iodine into hydrobromic and hydriodic acids in the presence of water. For example, in the action of phosphorus the essence of the matter is that the oxygen of the water goes to the phosphorus, and the union of the remaining elements leads to the formation of hydrobromic or hydriodic acid; but the matter is complicated by the reversibility of the reaction, the affinity for water, and other circumstances which are understood by following Berthollet's doctrine. Chlorine (and bromine also) directly decomposes hydrogen sulphide, forming hydrochloric acid and liberating sulphur, both in a gaseous form and in solutions, whilst iodine only decomposes hydrogen sulphide in weak solutions, when its affinity for hydrogen is aided by the affinity of hydrogen iodide for water. In a gaseous state iodine does not act on hydrogen sulphide,[73] whilst sulphur is able to decompose gaseous hydriodic acid, forming hydrogen sulphide and a compound of sulphur and iodine which with water forms hydriodic acid.[74]

[71] A number of similar cases confirm what has been said in Chapter X.

[72] This is prevented by the reducibility of sulphuric acid. If
volatile acids be taken they pass over, together with the
hydrobromic and hydriodic acids, when distilled; whilst many
non-volatile acids which are not reduced by hydrobromic and
hydriodic acids only act feebly (like phosphoric acid), or do not
act at all (like boric acid).

[73] This is in agreement with the thermochemical data, because if all
the substances be taken in the gaseous state (for sulphur the heat
of fusion is 0·3, and the heat of vaporisation 2·3) we have H_{2}
+ S = 4·7; H_{2} + Cl_{2} = 44; H_{2} + Br_{2} = 24, and H_{2} +
I_{2} = -3 thousand heat units; hence the formation of H_{2}S
gives less heat than that of HCl and HBr, but more than that of
HI. In dilute solutions H_{2} + S + Aq = 9·3, and consequently
less than the formation of all the halogen acids, as H_{2}S
evolves but little heat with water, and therefore in dilute
solutions chlorine, bromine, and iodine decompose hydrogen
sulphide.

[74] Here there are three elements, hydrogen, sulphur, and iodine, each
pair of which is able to form a compound, HI, H_{2}S, and SI,
besides which the latter may unite in various proportions. The
complexity of chemical mechanics is seen in such examples as
these. It is evident that only the study of the simplest cases can
give the key to the more complex problems, and on the other hand
it is evident from the examples cited in the last pages that,
without penetrating into the conditions of chemical equilibria, it
would be impossible to explain chemical phenomena. By following
the footsteps of Berthollet the possibility of unravelling the
problems will be reached; but work in this direction has only been
begun during the last ten years, and much remains to be done in
collecting experimental material, for which occasions present
themselves at every step. In speaking of the halogens I wished to
turn the reader's attention to problems of this kind.

If hydrogen sulphide be passed through water containing iodine, the reaction H_{2}S + I_{2} = 2HI + S proceeds so long as the solution is dilute, but when the mass of free HI increases the reaction stops, because the iodine then passes into solution. A solution having a composition approximating to 2HI + 4I_{2} + 9H_{2}O (according to Bineau) does not react with H_{2}S, notwithstanding the quantity of free iodine. Therefore only weak solutions of hydriodic acid can be obtained by passing hydrogen sulphide into water with iodine.[74 bis]

[74 bis] The same essentially takes place when sulphurous anhydride,
in a dilute solution, gives hydriodic acid and sulphuric acid with
iodine. On concentration a reverse reaction takes place. The
equilibrated systems and the part played by water are everywhere
distinctly seen.

To obtain[75] gaseous hydrobromic and hydriodic acids it is most convenient to take advantage of the reactions between phosphorus, the halogens, and water, the latter being present in small quantity (otherwise the halogen acids formed are dissolved by it); the halogen is gradually added to the phosphorus moistened with water. Thus if red phosphorus be placed in a flask and moistened with water, and bromine be added drop by drop (from a tap funnel), hydrobromic acid is abundantly and uniformly disengaged.[76] Hydrogen iodide is prepared by adding 1 part of common (yellow) dry phosphorus to 10 parts of dry iodine in a glass flask. On shaking the flask, union proceeds quietly between them (light and heat being evolved), and when the mass of iodide of phosphorus which is formed has cooled, water is added drop by drop (from a tap funnel) and hydrogen iodide is evolved directly without the aid of heat. These methods of preparation will be at once understood when it is remembered (p. 468) that phosphorus chloride gives hydrogen chloride with water. It is exactly the same here--the oxygen of the water passes over to the phosphorus, and the hydrogen to the iodine, thus, PI_{3} + 3H_{2}O = PH_{3}O_{3} + 3HI.[77]

[75] Methods of formation and preparation are nothing more than
particular cases of chemical reaction. If the knowledge of
chemical mechanics were more exact and complete than it now is it
would be possible to foretell all cases of preparation _with every
detail_ (of the quantity of water, temperature, pressure, mass,
&c.) The study of practical methods of preparation is therefore
one of the paths for the study of chemical mechanics. The reaction
of iodine on phosphorus and water is a case like that mentioned in
Note 74, and the matter is here further complicated by the
possibility of the formation of the compound PH_{3} with HI, as
well as the production of PI_{2}, PI_{3}, and the affinity of
hydriodic acid and the acids of phosphorus for water. The
theoretical interest of equilibria in all their complexity is
naturally very great, but it falls into the background in presence
of the primary interest of discovering practical methods for the
isolation of substances, and the means of employing them for the
requirements of man. It is only after the satisfaction of these
requirements that interests of the other order arise, which in
their turn must exert an influence on the former. For these
reasons, whilst considering it opportune to point out the
theoretical interest of chemical equilibria, the chief attention
of the reader is directed in this work to questions of practical
importance.

[76] Hydrobromic acid is also obtained by the action of bromine on
paraffin heated to 180°. Gustavson proposed to prepare it by the
action of bromine (best added in drops together with traces of
aluminium bromide) on anthracene (a solid hydrocarbon from coal
tar). Balard prepared it by passing bromine vapour over moist
pieces of common phosphorus. The liquid tribromide of phosphorus,
directly obtained from phosphorus and bromine, also gives
hydrobromic acid when treated with water. Bromide of potassium or
sodium, when treated with sulphuric acid in the presence of
phosphorus, also gives hydrobromic acid, but hydriodic acid is
decomposed by this method. In order to free hydrobromic acid from
bromine vapour it is passed over moist phosphorus and dried either
by phosphoric anhydride or calcium bromide (calcium chloride
cannot be used, as hydrochloric acid would be formed). Neither
hydrobromic nor hydriodic acids can be collected over mercury, on
which they act, but they may be directly collected in a dry vessel
by leading the gas-conducting tube to the bottom of the vessel,
both gases being much heavier than air. Merz and Holtzmann (1889)
proposed to prepare HBr directly from bromine and hydrogen. For
this purpose pure dry hydrogen is passed through a flask
containing boiling bromine. The mixture of gas and vapour then
passes through a tube provided with one or two bulbs, which is
heated moderately in the middle. Hydrobromic acid is formed with a
series of flashes at the part heated. The resultant HBr, together
with traces of bromine, passes into a Woulfe's bottle into which
hydrogen is also introduced, and the mixture is then carried
through another heated tube, after which it is passed through
water which dissolves the hydrobromic acid. According to the
method proposed by Newth (1892) a mixture of bromine and hydrogen
is led through a tube containing a platinum spiral, which is
heated to redness after the air has been displaced from the tube.
If the vessel containing the bromine be kept at 60°, the hydrogen
takes up almost the theoretical amount of bromine required for the
formation of HBr. Although the flame which appears in the
neighbourhood of the platinum spiral does not penetrate into the
vessel containing the bromine, still, for safety, a tube filled
with cotton wool may be interposed.

Hydroiodic acid is obtained in the same manner as hydrobromic. The
iodine is heated in a small flask, and its vapour is carried over
by hydrogen into a strongly heated tube, The gas passing from the
tube is found to contain a considerable amount of HI, together
with some free iodine. At a low red heat about 17 p.c. of the
iodine vapour enters into combination; at a higher temperature, 78
p.c. to 79 p.c.; and at a strong heat about 82 p.c.

[77] But generally more phosphorus is taken than is required for the
formation of PI_{3}, because otherwise a portion of the iodine
distils over. If less than one-tenth part of iodine be taken, much
phosphonium iodide, PH_{4}I, is formed. This proportion was
established by Gay-Lussac and Kolbe. Hydriodic acid is also
prepared in many other ways. Bannoff dissolves two parts of iodine
in one part of a previously prepared strong (sp. gr. 1·67)
solution of hydriodic acid, and pours it on to red phosphorus in a
retort. Personne takes a mixture of fifteen parts of water, ten of
iodine, and one of red phosphorus, which, when heated, disengages
hydriodic acid mixed with iodine vapour; the latter is removed by
passing it over moist phosphorus (Note 76). It must be remembered
however that reverse reaction (Oppenheim) may take place between
the hydriodic acid and phosphorus, in which the compounds PH_{4}I
and PI_{2} are formed.

It should be observed that the reaction between phosphorus, iodine
and water must be carried out in the above proportions and with
caution, as they may react with explosion. With red phosphorus the
reaction proceeds quietly, but nevertheless requires care.

L. Meyer showed that with an excess of iodine the reaction
proceeds without the formation of bye-products (PH_{4}I),
according to the equation P + 5I + 4H_{2}O = PH_{3}O_{4} + 5HI.
For this purpose 100 grams of iodine and 10 grams of water are
placed in a retort, and a paste of 5 grams of red phosphorus and
10 grams of water is added little by little (at first with great
care). The hydriodic acid may be obtained free from iodine by
directing the neck of the retort upwards and causing the gas to
pass through a shallow layer of water (respecting the formation of
HI, _see_ also Note 75).

In a gaseous form hydrobromic and hydriodic acids are closely analogous to hydrochloric acid; they are liquefied by pressure and cold, they fume in the air, form solutions and hydrates, of constant boiling point, and react on metals, oxides and salts, &c.[78] Only the relatively easy decomposability of hydrobromic acid and especially of hydriodic acid, clearly distinguish these acids from hydrochloric acid. For this reason, hydriodic acid acts in a number of cases as a deoxidiser or reducer, and frequently even serves as a means for the transference of hydrogen. Thus Berthelot, Baeyer, Wreden, and others, by heating unsaturated hydrocarbons in a solution of hydriodic acid, obtained their compounds with hydrogen nearer to the limit C_{_n_}H_{2_n_ + 2} or even the saturated compounds. For example, benzene, C_{6}H_{6}, when heated in a closed tube with a strong solution of hydriodic acid, gives hexylene, C_{6}H_{12}. The easy decomposability of hydriodic acid accounts for the fact that iodine does not act by metalepsis on hydrocarbons, for the hydrogen iodide liberated with the product of metalepsis, RI, formed, gives iodine and the hydrogen compound, RH, back again. And therefore, to obtain the products of iodine substitution, either iodic acid, HIO_{3} (Kekulé), or mercury oxide, HgO (Weselsky), is added, as they immediately react on the hydrogen iodide, thus: HIO_{3} + 5HI = 3H_{2}O + 3I_{2}, or, HgO + 2HI = HgI_{2} + H_{2}O. From these considerations it will be readily understood that iodine acts like chlorine (or bromine) on ammonia and sodium hydroxide, for in these cases the hydriodic acid produced forms NH_{4}I and NaI. With tincture of iodine or even the solid element, a solution of ammonia immediately forms a highly-explosive solid black product of metalepsis, NHI_{2}, generally known as _iodide_ of _nitrogen_, although it still contains hydrogen (this was proved beyond doubt by Szuhay 1893), which may be replaced by silver (with the formation of NAgI_{2}): 3NH_{3} + 2I_{2} = 2NH_{4}I + NHI_{2}. However, the composition of the last product is variable, and with an excess of water NI_{3} seems to be formed. Iodide of nitrogen is just as explosive as nitrogen chloride.[78 bis] In the action of iodine on sodium hydroxide no bleaching compound is formed (whilst bromine gives one), but a direct reaction is always accomplished with the formation of an iodate, 6NaHO + 3I_{2} = 5NaI + 3H_{2}O + NaIO_{3} (Gay-Lussac). Solutions of other alkalis, and even a mixture of water and oxide of mercury, act in the same manner.[79] This direct formation of _iodic acid_, HIO_{3} = IO_{2}(OH), shows the propensity of iodine to give compounds of the type IX_{5}. Indeed, this capacity of iodine to form compounds of a high type emphasises itself in many ways. But it is most important to turn attention to the fact that iodic acid is easily and directly formed by the action of oxidising substances on iodine. Thus, for instance, strong nitric acid directly converts iodine into iodic acid, whilst it has no oxidising action on chlorine.[79 bis] This shows a greater affinity in iodine for oxygen than in chlorine, and this conclusion is confirmed by the fact that iodine displaces chlorine from its oxygen acids,[80] and that in the presence of water chlorine oxidises iodine.[81] Even ozone or a silent discharge passed through a mixture of oxygen and iodine vapour is able to directly oxidise iodine[82] into iodic acid. It is disengaged from solutions as a hydrate, HIO_{3}, which loses water at 170°, and gives an anhydride, I_{2}O_{5}. Both these substances are crystalline (sp. gr. I_{2}O_{5} 5·037, HIO_{3} 4·869 at 0°), colourless and soluble in water;[83] both decompose at a red heat into iodine and oxygen, are in many cases powerfully oxidising--for instance, they oxidise sulphurous anhydride, hydrogen sulphide, carbonic oxide, &c.--form chloride of iodine and water with hydrochloric acid, and with bases form salts, not only normal MIO_{3}, but also acid; for example, KIO_{3}HIO_{3}, KIO_{3}2HIO_{3}.[83 bis] With hydriodic acid iodic acid immediately reacts, disengaging iodine, HIO_{3} + 5HI = 3H_{2}O + 3I_{2}.

[78] The specific gravities of their solutions as deduced by me on the
basis of Topsöe and Berthelot's determinations for 15°/4° are as
follows:--

10 20 30 40 50 60 p.c.
HBr 1·071 1·156 1·258 1·374 1·505 1·650
HI 1·075 1·164 1·267 1·399 1·567 1·769

Hydrobromic acid forms two hydrates, HBr,2H_{2}O and HBr,H_{2}O,
which have been studied by Roozeboom with as much completeness as
the hydrate of hydrochloric acid (Chapter X. Note 37).

With metallic silver, solutions of hydriodic acid give hydrogen
with great ease, forming silver iodide. Mercury, lead, and other
metals act in a similar manner.

[78 bis] Iodide of nitrogen, NHI_{2} is obtained as a brown pulverulent
precipitate on adding a solution of iodine (in alcohol, for
instance) to a solution of ammonia. If it be collected on a
filter-paper, it does not decompose so long as the precipitate is
moist; but when dry it explodes violently, so that it can only be
experimented upon in small quantities. Usually the filter-paper is
torn into bits while moist, and the pieces laid upon a brick; on
drying an explosion proceeds not only from friction or a blow, but
even spontaneously. The more dilute the solution of ammonia, the
greater is the amount of iodine required for the formation of the
precipitate of NHI_{2}. A low temperature facilitates its
formation. NHI_{2} dissolves in ammonia water, and when heated the
solution forms HIO_{3} and iodine. With KI, iodide of nitrogen
gives iodine, NH_{3} and KHO. These reactions (Selivanoff) are
explained by the formation of HIO from NHI_{2} + 2H_{2}O = NH_{3} +
2HIO--and then KI + HIO = I_{2} + KHO. Selivanoff (_see_ Note 29)
usually observed a temporary formation of hypoiodous acid, HIO, in
the reaction of ammonia upon iodine, so that here the formation of
NHI_{2} is preceded by that of HIO--_i.e._ first I_{2} + H_{2}O =
HIO + HI, and then not only the HI combines with NH_{3}, but also
2HIO + NH_{3} = NHI_{2} + 2H_{2}O. With dilute sulphuric acid
iodide of nitrogen (like NCl_{3}) forms hypoiodous acid, but it
immediately passes into iodic acid, as is expressed by the
equation 5HIO = 2I_{2} + HIO_{3} + 2H_{2}O (first 3HIO = HIO_{3} +
2HI, and then HI + HIO = I_{2} + H_{2}O). Moreover, Selivanoff
found that iodide of nitrogen, NHI_{2}, dissolves in an excess of
ammonia water, and that with potassium iodide the solution gives
the reaction for hypoiodous acid (the evolution of iodine in an
alkaline solution). This shows that HIO participates in the
formation and decomposition of NHI_{2}, and therefore the
condition of the iodine (its metaleptic position) in them is
analogous, and differs from the condition of the halogens in the
haloid-anhydrides (for instance, NO_{2}Cl). The latter are
tolerably stable, while (the haloid being designated by X)
NHX_{2}, NX_{3}, XOH, RXO (_see_ Chapter XIII. Note 43), &c., are
unstable, easily decomposed with the evolution of heat, and, under
the action of water, the haloid is easily replaced by hydrogen
(Selivanoff), as would be expected in true products of metalepsis.

[79] Hypoiodous acid, HIO, is not known, but organic compounds, RIO, of
this type are known. To illustrate the peculiarities of their
properties we will mention one of these compounds, namely,
_iodosobenzol_, C_{6}H_{5}IO. This substance was obtained by
Willgerodt (1892), and also by V. Meyer, Wachter, and Askenasy, by
the action of caustic alkalis upon phenoldiiodochloride,
C_{6}H_{5}ICl_{2} (according to the equation, C_{6}H_{5}ICl_{2} +
2MOH = C_{6}H_{5}IO + 2MCl + H_{2}O). Iodosobenzol is an amorphous
yellow substance, whose melting point could not be determined
because it explodes at 210°, decomposing with the evolution of
iodine vapour. This substance dissolves in hot water and alcohol,
but is not soluble in the majority of other neutral organic
solvents. If acids do not oxidise C_{6}H_{5}IO, they give saline
compounds in which iodosobenzol appears as a basic oxide of a
diatomic metal, C_{6}H_{5}I. Thus, for instance, when an acetic
acid solution of iodosobenzol is treated with a solution of nitric
acid, it gives large monoclinic crystals of a nitric acid salt
having the composition C_{6}H_{5}I(NO_{3})_{2} (like
Ca(NO_{3})_{2}). In appearing as the analogue of basic oxides,
iodosobenzol displaces iodine from potassium iodide (in a solution
acidulated with acetic or hydrochloric acid)--_i.e._ it acts with
its oxygen like HClO. The action of peroxide of hydrogen, chromic
acid, and other similar oxidising agents gives iodoxybenzol,
C_{6}H_{5}IO_{2}, which is a neutral substance--_i.e._ incapable
of giving salts with acids (compare Chapter XIII. Note 43).

[79 bis] The oxidation of iodine by strong nitric acid was discovered
by Connell; Millon showed that it is effected, although more
slowly, by the action of the hydrates of nitric acid up to
HNO_{3},H_{2}O, but that the solution HNO_{3},2H_{2}O, and weaker
solutions, do not oxidise, but simply dissolve, iodine. The
participation of water in reactions is seen in this instance. It
is also seen, for example, in the fact that dry ammonia combines
directly with iodine--for instance, at 0° forming the compound
I_{2},4NH_{3}--whilst iodide of nitrogen is only formed in
presence of water.

[80] Bromine also displaces chlorine--for instance, from chloric acid,
directly forming bromic acid. If a solution of potassium chlorate
be taken (75 parts per 400 parts of water), and iodine be added to
it (80 parts), and then a small quantity of nitric acid, chlorine
is disengaged on boiling, and potassium iodate is formed in the
solution. In this instance the nitric acid first evolves a certain
portion of the chloric acid, and the latter, with the iodine,
evolves chlorine. The iodic acid thus formed acts on a further
quantity of the potassium chlorate, sets a portion of the chloric
acid free, and in this manner the action is kept up. Potilitzin
(1887) remarked, however, that not only do bromine and iodine
displace the chlorine from chloric acid and potassium chlorate,
but also chlorine displaces bromine from sodium bromate, and,
furthermore, the reaction does not proceed as a direct
substitution of the halogens, but is accompanied by the formation
of free acids; for example, 5NaClO_{3} + 3Br_{2} + 3H_{2}O = 5NaBr
+ 5HClO_{3} + HBrO_{3}.

[81] If iodine be stirred up in water, and chlorine passed through the
mixture, the iodine is dissolved; the liquid becomes colourless,
and contains, according to the relative amounts of water and
chlorine, either IHCl_{2}, or ICl_{3}, or HIO_{3}. If there be a
small amount of water, then the iodic acid may separate out
directly as crystals, but a complete conversion (Bornemann) only
occurs when not less than ten parts of water are taken to one part
of iodine--ICl + 3H_{2}O + 2Cl_{2} = IHO_{3} + 5HCl.

[82] Schönbein and Ogier proved this. Ogier found that at 45° ozone
immediately oxidises iodine vapour, forming first of all the oxide
I_{2}O_{3}, which is decomposed by water or on heating into iodic
anhydride and iodine. Iodic acid is formed at the positive pole
when a solution of hydriodic acid is decomposed by a galvanic
current (Riche). It is also formed in the combustion of hydrogen
mixed with a small quantity of hydriodic acid (Salet).

[83] Kämmerer showed that a solution of sp. gr. 2·127 at 14°,
containing 2HIO_{3},9H_{2}O, solidified completely in the cold. On
comparing solutions HI + _m_H_{2}O with HIO_{3} + _m_H_{2}O, we
find that the specific gravity increases but the volume decreases,
whilst in the passage of solutions HCl + _m_H_{2}O to HClO_{3} +
_m_H_{2}O both the specific gravity and the volume increase, which
is also observed in certain other cases (for example, H_{3}PO_{3}
and H_{3}PO_{4}).

[83 bis] Ditte (1890) obtained many iodates of great variety. A neutral
salt, 2(LiIO_{3})H_{2}O, is obtained by saturating a solution of
lithia with iodic acid. There is an analogous ammonium salt,
2(NH_{4}IO_{3})H_{2}O. He also obtained hydrates of a more complex
composition, such as 6(NH_{4}IO_{3})H_{2}O and
6(NH_{4}IO_{3})2H_{2}O. Salts of the alkaline earths,
Ba(IO_{3})_{2}H_{2}O and Sr(IO_{3})_{2}H_{2}O, may be obtained by
a reaction of double decomposition from the normal salts of the
type 2(MeIO_{3})H_{2}O. When evaporated at 70° to 80° with nitric
acid these salts lose water. A mixture of solutions of nitrate of
zinc and an alkaline iodate precipitates Zn(IO_{3})_{2}2H_{2}O. An
anhydrous salt is thrown out if nitric acid be added to the
solutions. Analogous salts of cadmium, silver, and copper give
compounds of the type 2Me´IO_{3}4NH_{3} and
Me´´(IO_{3})_{2}4NH_{3}, with gaseous ammonia (Me´ and Me´´ being
elements of the first (Ag) and second (Cd, Zn, Cu) groups). With
an aqueous solution of ammonia the above salts give substances of
a different composition, such as Zn(IO_{3})_{2}(NH_{4})_{2}O,
Cd(IO_{3})_{2}(NH_{4})_{2}O. Copper gives
Cu(IO_{3})_{2}4(NH_{4})_{2}O and Cu(IO_{3})_{2}(NH_{4})_{2}O.
These salts may be regarded as compounds of I_{2}O_{5}, and MeO
and (NH_{4})_{2}O; for example, Zn(IO_{3})_{2}(NH_{4})_{2}O may be
regarded as ZnO(NH_{4})_{2}OI_{2}O_{5}, or, as derived from the
hydrate, I_{2}O_{5}2H_{2}O = 2(HIO_{3})H_{2}O.

As with chlorine, so with iodine, a _periodic acid_, HIO_{4}, is formed. This acid is produced in the form of its salts, by the action of chlorine on alkaline solutions of iodates, and also by the action of iodine on chloric acid.[84] It crystallises from solutions as a hydrate containing 2H_{2}O (corresponding with HClO_{4},2H_{2}O), but as it forms salts containing up to 5 atoms of metals, this water must be counted as water of constitution. Therefore IO(OH)_{5} = HIO_{4},2H_{2}O corresponds with the highest form of halogen compounds, IX_{7}.[85] In decomposing (at 200°) or acting as an oxidiser, periodic acid first gives iodic acid, but it may also be ultimately decomposed.

[84] If sodium iodate be mixed with a solution of sodium hydroxide,
heated, and chlorine passed through the solution, a sparingly
soluble salt separates out, which corresponds with periodic acid,
and has the composition Na_{4}I_{2}O_{9},3H_{2}O.

6NaHO + 2NaIO_{3} + 4Cl = 4NaCl + Na_{4}I_{2}O_{9} + 3H_{2}O.

This compound is sparingly soluble in water, but dissolves easily
in a very dilute solution of nitric acid. If silver nitrate be
added to this solution a precipitate is formed which contains the
corresponding compound of silver, Ag_{4}I_{2}O_{9},3H_{2}O. If
this sparingly soluble silver compound be dissolved in hot nitric
acid, orange crystals of a salt having the composition AgIO_{4}
separate on evaporation. This salt is formed from the preceding by
the nitric acid taking up silver oxide--Ag_{4}I_{2}O_{9} +
2HNO_{3} = 2AgNO_3 + 2AgIO_{4} + H_{2}O. The silver salt is
decomposed by water, with the re-formation of the preceding salt,
whilst iodic acid remains in solution--

4AgIO_{4} + H_{2}O = Ag_{4}I_{2}O_{9} + 2HIO_{4}.

The structure of the first of these salts,
Na_{4}I_{2}O_{9},3H_{2}O, presents itself in a simpler form if the
water of crystallisation is regarded as an integral portion of the
salt; the formula is then divided in two, and takes the form of
IO(OH)_{3}(ONa)_{2}--that is, it answers to the type IOX_{5}, or
IX_{7}, like AgIO_{4} which is IO_{3}(OAg). The composition of all
the salts of periodic acids are expressed by this type IX_{7}.
Kimmins (1889) refers all the salts of periodic acid to four
types--the meta-salts of HIO_{4} (salts of Ag, Cu, Pb), the
meso-salts of H_{3}IO_{5} (PbH, Ag_{2}H, CdH), the para-salts of
H_{5}IO_{6} (Na_{2}H_{3}, Na_{3}H_{2}), and the di-salts of
H_{4}I_{2}O_{9} (K_{4}, Ag_{4}, Ni_{2}). The three first are
direct compounds of the type IX_{7}, namely, IO_{3}(OH),
IO_{2}(OH)_{3}, and IO(OH)_{5}, and the last are types of
diperiodic salts, which correspond with the type of the
meso-salts, as pyrophosphoric salts correspond with
orthophosphoric salts--_i.e._ 2H_{3}IO_{5}-H_{2}O =
H_{4}I_{2}O_{9}.

[85] Periodic acid, discovered by Magnus and Ammermüller, and whose
salts were afterwards studied by Langlois, Rammelsberg, and many
others, presents an example of hydrates in which it is evident
that there is not that distinction between the water of hydration
and of crystallisation which was at first considered to be so
clear. In HClO,2H_{2}O the water, 2H_{2}O, is not displaced by
bases, and must be regarded as water of crystallisation, whilst in
HIO_{4},2H_{2}O it must be regarded as water of hydration. We
shall afterwards see that the system of the elements obliges us to
consider the halogens as substances giving a highest saline type,
_GX__{7}, where _G_ signifies a halogen, and _X_ oxygen (O =
_X__{2}), OH, and other like elements. The hydrate IO(OH)_{5}
corresponding with many of the salts of periodic acid (for
example, the salts of barium, strontium, mercury) does not exhaust
all the possible forms. It is evident that various other pyro-,
meta-, &c., forms are possible by the loss of water, as will be
more fully explained in speaking of phosphoric acid, and as was
pointed out in the preceding note.

Compounds formed between chlorine and iodine must be classed among the most interesting halogen bodies.[86] These elements combine together directly with evolution of heat, and form _iodine monochloride_, ICl, or _iodine trichloride_, ICl_{3}.[87] As water reacts on these substances, forming iodic acid and iodine, they have to be prepared from dry iodine and chlorine.[88] Both substances are formed in a number of reactions; for example, by the action of aqua regia on iodine, of chlorine on hydriodic acid, of hydrochloric acid on periodic acid, of iodine on potassium chlorate (with the aid of heat, &c.) Trapp obtained iodine monochloride, in beautiful red crystals, by passing a rapid current of chlorine into molten iodine. The monochloride then distils over and solidifies, melting at 27°. By passing chlorine over the crystals of the monochloride, it is easy to obtain iodine trichloride in orange crystals, which melt at 34° and volatilise at 47°, but in so doing decompose (into Cl_{2} and ClI). The chemical properties of these chlorides entirely resemble those of chlorine and iodine, as would be expected, because, in this instance, a combination of similar substances has taken place as in the formation of solutions or alloys. Thus, for instance, the unsaturated hydrocarbons (for example, C_{2}H_{4}), which are capable of directly combining with chlorine and iodine, also directly combine with iodine monochloride.

[86] With respect to hydrogen, oxygen, chlorine, and other elements,
bromine occupies an intermediate position between chlorine and
iodine, and therefore there is no particular need for considering
at length the compounds of bromine. This is the great advantage of
a natural grouping of the elements.

[87] They were both obtained by Gay-Lussac and many others. Recent data
respecting iodine monochloride, ICl, entirely confirm the numerous
observations of Trapp (1854), and even confirm his statement as to
the existence of two isomeric (liquid and crystalline) forms
(Stortenbeker). With a small excess of iodine, iodine monochloride
remains liquid, but in the presence of traces of iodine
trichloride it easily crystallises. Tanatar (1893) showed that of
the two modifications of ICl, one is stable, and melts at 27°;
while the other, which easily passes into the first, and is formed
in the absence of ICl_{3}, melts at 14°. Schützenberger amplified
the data concerning the action of water on the chlorides (Note
88), and Christomanos gave the fullest data regarding the
trichloride.

After being kept for some time, the liquid monochloride of iodine
yields red deliquescent octahedra, having the composition ICl_{4},
which are therefore formed from the monochloride with the
liberation of free iodine, which dissolves in the remaining
quantity of the monochloride. This substance, however, judging by
certain observations, is impure iodine trichloride. If 1 part of
iodine be stirred up in 20 parts of water, and chlorine be passed
through the liquid, then all the iodine is dissolved, and a
colourless liquid is ultimately obtained which contains a certain
proportion of chlorine, because this compound gives a metallic
chloride and iodate with alkalis without evolving any free iodine:
ICl_{5} + 6KHO = 5KCl + KIO_{3} + 3H_{2}O. The existence of a
pentachloride ICl_{5} is, however, denied, because this substance
has not been obtained in a free state.

Stortenbeker (1888) investigated the equilibrium of the system
containing the molecules I_{2}, ICl, ICl_{3}, and Cl_{2}, in the
same way that Roozeboom (Chapter X. Note 38) examined the
equilibrium of the molecules HCl, HCl,2H_{2}O, and H_{2}O. He
found that iodine monochloride appears in two states, one (the
ordinary) is stable and melts at 27°·2, whilst the other is
obtained by rapid cooling, and melts at 13°·9, and easily passes
into the first form. Iodine trichloride melts at 101° only in a
closed tube under a pressure of 16 atmospheres.

[88] By the action of water on iodine monochloride and trichloride a
compound IHCl_{2} is obtained, which does not seem to be altered
by water. Besides this compound, iodine and iodic acid are always
formed, 10ICl + 3H_{2}O = HIO_{3} + 5IHCl_{2} + 2I_{2}; and in
this respect iodine trichloride may be regarded as a mixture, ICl
+ ICl_{5} = 2ICl_{3}, but ICl_{5} + 3H_{2}O = IHO_{3} + 5HCl;
hence iodic acid, iodine, the compound IHCl_{2}, and hydrochloric
acid are also formed by the action of water.

CHAPTERR XII

SODIUM

The neutral salt, sodium sulphate, Na_{2}SO_{4}, obtained when a mixture of sulphuric acid and common salt is strongly heated (Chapter X.),[1] forms a colourless saline mass consisting of fine crystals, soluble in water. It is the product of many other double decompositions, sometimes carried out on a large scale; for example, when ammonium sulphate is heated with common salt, in which case the sal-ammoniac is volatilised, &c. A similar decomposition also takes place when, for instance, a mixture of lead sulphate and common salt is heated; this mixture easily fuses, and if the temperature be further raised heavy vapours of lead chloride appear. When the disengagement of these vapours ceases, the remaining mass, on being treated with water, yields a solution of sodium sulphate mixed with a solution of undecomposed common salt. A considerable quantity, however, of the lead sulphate remains unchanged during this reaction, PbSO_{4} + 2NaCl = PbCl_{2} + Na_{2}SO_{4}, the vapours will contain lead chloride, and the residue will contain the mixture of the three remaining salts. The cause and nature of the reaction are just the same as were pointed out when considering the action of sulphuric acid upon NaCl. Here too it may be shown that the double decomposition is determined by the removal of PbCl_{2} from the sphere of the action of the remaining substances. This is seen from the fact that sodium sulphate, on being dissolved in water and mixed with a solution of any lead salt (and even with a solution of lead chloride, although this latter is but sparingly soluble in water), immediately gives a white precipitate of lead sulphate. In this case the lead takes up the elements of sulphuric acid from the sodium sulphate in the solutions. On heating, the reverse phenomenon is observed. The reaction in the solution depends upon the insolubility of the lead sulphate, and the decomposition which takes place on heating is due to the volatility of the lead chloride. Silver sulphate, Ag_{2}SO_{4}, in solution with common salt, gives silver chloride, because the latter is insoluble in water, Ag_{2}SO_{4} + 2NaCl = Na_{2}SO_{4} + 2AgCl. Sodium carbonate, mixed in solution with the sulphates of iron, copper, manganese, magnesium, &c., gives in solution sodium sulphate, and in the precipitate a carbonate of the corresponding metal, because these salts of carbonic acid are insoluble in water; for instance, MgSO_{4} + Na_{2}CO_{3} = Na_{2}SO_{4} + MgCO_{3}. In precisely the same way sodium hydroxide acts on solutions of the majority of the salts of sulphuric acid containing metals, the hydroxides of which are insoluble in water--for instance, CuSO_{4} + 2NaHO = Cu(HO)_{2} + Na_{2}SO_{4}. Sulphate of magnesium, MgSO_{4}, on being mixed in solution with common salt, forms, although not completely, chloride of magnesium, and sodium sulphate. On cooling the mixture of such (concentrated) solutions sodium sulphate is deposited, as was shown in Chapter X. This is made use of for preparing it on the large scale in works where sea-water is treated. In this case, on cooling, the reaction 2NaCl + MgSO_{4} = MgCl_{2} + Na_{2}SO_{4} takes place.

[1] Whilst describing in some detail the properties of sodium
chloride, hydrochloric acid, and sodium sulphate, I wish to impart,
by separate examples, an idea of the properties of saline
substances, but the dimensions of this work and its purpose and aim
do not permit of entering into particulars concerning every salt,
acid, or other substance. The fundamental object of this work--an
account of the characteristics of the elements and an acquaintance
with the forces acting between atoms--has nothing to gain from the
multiplication of the number of as yet ungeneralised properties and
relations.