chlorine when cold, but at a red heat, in a current of chlorine,
it forms calcium chloride, with the evolution of oxygen. (This was
confirmed in 1893 by Wells, at Oxford.) This reaction corresponds
with the decomposing action of chlorine on methane, ammonia, and
water. Slaked lime (calcium hydroxide, CaH_{2}O_{2}) also, when
dry, does not absorb chlorine at 100°. The absorption proceeds at
the ordinary temperature (below 40°). The dry mass thus obtained
contains not less than three equivalents of calcium hydroxide to
four equivalents of chlorine, so that its composition is
[Ca(HO)_{2}]_{5}Cl_{4}. In all probability a simple absorption of
chlorine by the lime at first takes place in this case, as may be
seen from the fact that even carbonic anhydride, when acting on
the dry mass obtained as above, disengages all the chlorine from
it, leaving only calcium carbonate. But if the bleaching powder be
obtained by a wet method, or if it be dissolved in water (in which
it is very soluble), and carbonic anhydride be passed into it,
then chlorine is no longer disengaged, but chlorine oxide,
Cl_{2}O, and only half of the chlorine is converted into this
oxide, while the other half remains in the liquid as calcium
chloride. From this it may be inferred that calcium chloride is
formed by the action of water on bleaching powder, and this is
proved to be the case by the fact that small quantities of water
extract a considerable amount of calcium chloride from bleaching
powder. If a large quantity of water act on bleaching powder an
excess of calcium hydroxide remains, a portion of which is not
subjected to change. The action of the water may be expressed by
the following formulæ: From the dry mass Ca_{3}(HO)_{6}Cl_{4}
there is formed lime, Ca(HO)_{2}, calcium chloride, CaCl_{2}, and
a saline substance, Ca(ClO)_{2}. Ca_{3}H_{6}O_{6}Cl_{4} =
CaH_{2}O_{2} + CaCl_{2}O_{2} + CaCl_{2} + 2H_{2}O. The resulting
substances are not equally soluble; water first extracts the
calcium chloride, which is the most soluble, then the compound
Ca(ClO)_{2} and ultimately calcium hydroxide is left. A mixture of
calcium chloride and hypochlorite passes into solution. On
evaporation there remains Ca_{2}O_{2}Cl_{4}3H_{2}O. The dry
bleaching powder does not absorb more chlorine, but the solution
is able to absorb it in considerable quantity. If the liquid be
boiled, a considerable amount of chlorine monoxide is evolved.
After this calcium chloride alone remains in solution, and the
decomposition may be expressed as follows: CaCl_{2} +
CaCl_{2}O_{2} + 2Cl_{2} = 2CaCl_{2} + 2Cl_{2}O. Chlorine monoxide
may be prepared in this manner.
It is sometimes said that bleaching powder contains a substance,
Ca(OH)_{2}Cl_{2}, that is calcium peroxide, CaO_{2}, in which one
atom of oxygen is replaced by (OH)_{2}, and the other by Cl_{2};
but, judging from what has been said above, this can only be the
case in the dry state, and not in solutions.
On being kept for some time, bleaching powder sometimes
decomposes, with the evolution of oxygen (because CaCl_{2}O_{2} =
CaCl_{2} + O_{2}, _see_ p. 163); the same takes place when it is
heated.
The products of the metalepsis of alkaline hydrates, NaClO and
Ca(ClO)_{2}, which are present in solutions of 'Javelle salt' and
bleaching powder (they are not obtained free from metallic chlorides),
must be counted as salts, because their metals are capable of
substitution. But the hydrate HClO corresponding with these salts,
or _hypochlorous acid_, is not obtained in a free or pure state, for
two reasons: in the first place, because this hydrate, as a very
feeble acid, splits up (like H_{2}CO_{3} or HNO_{3}) into water and the
anhydride, or _chlorine monoxide_, Cl_{2}O = 2HClO-H_{2}O; and, in the
second place, because, in a number of instances, it evolves oxygen
with great facility, forming hydrochloric acid: HClO = HCl + O. Both
hypochlorous acid and chlorine monoxide may be regarded as products of
the metalepsis of water, because HOH corresponds with ClOH and ClOCl.
Hence in many instances bleaching salts (a mixture of hypochlorites and
chlorides) break up, with the evolution of (1) _chlorine_, under the
action of an excess of a powerful acid capable of evolving hydrochloric
acid from sodium or calcium chlorides, and this takes place most simply
under the action of hydrochloric acid itself, because (p. 462) NaCl +
NaClO + 3HCl = 2NaCl + HCl + Cl_{2} + H_{2}O; (2) _oxygen_, as we saw
in Chapter III.--The bleaching properties and, in general, _oxidising
action_ of bleaching salts is based on this evolution of oxygen (or
chlorine); oxygen is also disengaged on heating the dry salts--for
instance, NaCl + NaClO = 2NaCl + O; (3) and, lastly, _chlorine
monoxide_, which contains both chlorine and oxygen. Thus, if a little
sulphuric, nitric, or similar acid (not enough to liberate hydrochloric
acid from the CaCl_{2}) be added to a solution of a bleaching salt
(which has an alkaline reaction, owing either to an excess of alkali
or to the feeble acid properties of HClO), then the hypochlorous acid
set free gives water and chlorine monoxide. If carbonic anhydride
(or boracic or a similar very feeble acid) act on the solution of a
bleaching salt, then hydrochloric acid is not evolved from the sodium
or calcium chlorides, but the hypochlorous acid is displaced and gives
chlorine monoxide,[31] because hypochlorous acid is one of the most
feeble acids. Another method for the preparation of chlorine monoxide
is based on these feeble acid properties of hypochlorous acid. Zinc
oxide and mercury oxide, under the action of chlorine in the presence
of water, do not give a salt of hypochlorous acid, but form a chloride
and hypochlorous acid, which fact shows the incapacity of this acid to
combine with the bases mentioned. Therefore, if such oxides as those of
zinc or mercury be shaken up in water, and chlorine be passed through
the turbid liquid,[32] a reaction occurs which may be expressed in the
following manner: 2HgO + 2Cl_{2} = Hg_{2}OCl_{2} + Cl_{2}O. In this case,
a compound of mercury oxide with mercury chloride, or the so-called
mercury oxychloride, is obtained: Hg_{2}OCl_{2} = HgO + HgCl_{2}. This is
insoluble in water, and is not affected by hypochlorous anhydride, so
that the solution will contain hypochlorous acid only, but the greater
part of it splits up into the anhydride and water.[32 bis]
[31] For this reason it is necessary that in the preparation of
bleaching powder the chlorine should be free from hydrochloric
acid, and even the lime from calcium chloride. An excess of
chlorine, in acting on a solution of bleaching powder, may also
give chlorine monoxide, because calcium carbonate also gives
chlorine monoxide under the action of chlorine. This reaction may
be brought about by treating freshly precipitated calcium
carbonate with a stream of chlorine in water: 2Cl_{2} + CaCO_{3} =
CO_{2} + CaCl_{2} + Cl_{2}O. From this we may conclude that,
although carbonic anhydride displaces hypochlorous anhydride, it
may be itself displaced by an excess of the latter.
[32] Dry red mercury oxide acts on chlorine, forming dry hypochlorous
anhydride (chlorine monoxide) (Balard); when mixed with water, red
mercury oxide acts feebly on chlorine, and when freshly
precipitated it evolves oxygen and chlorine. An oxide of mercury
which easily and abundantly evolves chlorine monoxide under the
action of chlorine in the presence of water may be prepared as
follows: the oxide of mercury, precipitated from a mercuric salt
by an alkali, is heated to 300° and cooled (Pelouze). If a salt,
MClO, be added to a solution of mercuric salt, HgX_{2}, mercuric
oxide is liberated, because the hypochlorite is decomposed.
[32 bis] A solution of hypochlorous anhydride is also obtained by the
action of chlorine on many salts; for example, in the action of
chlorine on a solution of sodium sulphate the following reaction
takes place: Na_{2}SO_{4} + H_{2}O + Cl_{2} = NaCl + HClO +
NaHSO_{4}. Here the hypochlorous acid is formed, together with
HCl, at the expense of chlorine and water, for Cl_{2} + H_{2}O =
HCl + HClO. If the crystallo-hydrate of chlorine be mixed with
mercury oxide, the hydrochloric acid formed in the reaction gives
mercury chloride, and hypochlorous acid remains in solution. A
dilute solution of hypochlorous acid or chlorine monoxide may be
concentrated by distillation, and if a substance which takes up
water (without destroying the acid)--for instance, calcium
nitrate--be added to the stronger solution, then the anhydride of
hypochlorous acid--_i.e._ chlorine monoxide--is disengaged.
Chlorine monoxide, which corresponds to bleaching and hypochlorous
salts, containing as it does the two elements oxygen and chlorine,
forms a characteristic example of a compound of elements which, in the
majority of cases, act chemically in an analogous manner. Chlorine
monoxide, as prepared from an aqueous solution by the abstraction of
water or by the action of dry chlorine on cold mercury oxide, is, at
the ordinary temperature, a gas or vapour which condenses into a red
liquid boiling at +20° and giving a vapour whose density (43 referred
to hydrogen) shows that 2 vols. of chlorine and 1 vol. of oxygen
give 2 vols. of chlorine monoxide. In an anhydrous form the gas or
liquid easily explodes, splitting up into chlorine and oxygen. This
explosiveness is determined by the fact that heat is _evolved_ in the
decomposition to the amount of about 15,000 heat units for Cl_{2}O.[33]
The explosion may even take place spontaneously, and also in the
presence of many oxidisable substances (for instance, sulphur, organic
compounds, &c.), but the solution, although unstable and showing a
strong oxidising tendency, does not explode.[34] It is evident that
the presence of hypochlorous acid, HClO, may be assumed in an aqueous
solution of Cl_{2}O, since Cl_{2}O + H_{2}O = 2HClO.
[33] All explosive substances are of this kind--ozone, hydrogen
peroxide, chloride of nitrogen, nitro-compounds, &c. Hence they
cannot be formed directly from the elements or their simplest
compounds, but, on the contrary, decompose into them. In a liquid
state chlorine monoxide explodes even on contact with powdery
substances, or when rapidly agitated--for instance, if a file be
rasped over the vessel in which it is contained.
[34] A solution of chlorine monoxide, or hypochlorous acid, does not
explode, owing to the presence of the mass of water. In
dissolving, chlorine monoxide evolves about 9,000 heat units, so
that its store of heat becomes less.
The capacity of hypochlorous acid (studied by Carius and others)
for entering into combination with the unsaturated hydrocarbons is
very often taken advantage of in organic chemistry. Thus its
solution absorbs ethylene, forming the chlorhydrin C_{2}H_{4}ClOH.
The oxidising action of hypochlorous acid and its salts is not
only applied to bleaching but also to many reactions of oxidation.
Thus it converts the lower oxides of manganese into the peroxide.
Hypochlorous acid, its salts, and chlorine monoxide serve as a transition
between hydrochloric acid, chlorides, and chlorine, and a whole series of
compounds containing the same elements combined with a still greater
quantity of oxygen. The higher oxides of chlorine, as their origin
indicates, are closely connected with hypochlorous acid and its salts:
Cl_{2}, NaCl, HCl, hydrochloric acid.
Cl_{2}O, NaClO, HClO, hypochlorous acid.
Cl_{2}O_{3}, NaClO_{2}, HClO_{2}, chlorous acid.[35]
Cl_{2}O_{5}, NaClO_{3}, HClO_{3}, chloric acid.
Cl_{2}O_{7}, NaClO_{4}, HClO_{4}, perchloric acid.
When heated, solutions of hypochlorites undergo a remarkable change.
Themselves so unstable, they, without any further addition, yield two
fresh salts which are both much more stable; one contains more oxygen
than MClO, the other contains none at all.
3MClO = MClO_{3} + 2MCl
hypochlorite chlorate chloride
[35] _Chlorous acid_, HClO_{2} (according to the data given by Millon,
Brandau, and others) in many respects resembles hypochlorous acid,
HClO, whilst they both differ from chloric and perchloric acids in
their degree of stability, which is expressed, for instance, in
their bleaching properties; the two higher acids do not bleach,
but both the lower ones do so (oxidise at the ordinary
temperature). On the other hand, chlorous acid is analogous to
nitrous acid, HNO_{2}. The anhydride of chlorous acid,
Cl_{2}O_{3}, is not known in a pure state, but it probably occurs
in admixture with chlorine dioxide, ClO_{2}, which is obtained by
the action of nitric and sulphuric acids on a mixture of potassium
chlorate with such reducing substances as nitric oxide, arsenious
oxide, sugar, &c. All that is at present known is that pure
chlorine dioxide ClO_{2} (_see_ Notes 39-43) is gradually
converted into a mixture of hypochlorous and chlorous acids under
the action of water (and alkalis); that is, it acts like nitric
peroxide, NO_{2} (giving HNO_{3} and HNO_{2}), or as a mixed
anhydride, 2ClO_{2} + H_{2}O = HClO_{3} + HClO_{2}. The silver
salt, AgClO_{2}, is sparingly soluble in water. The investigations
of Garzarolli-Thurnlackh and others seem to show that the
anhydride Cl_{2}O_{3} does not exist in a free state.
Part of the salt--namely, two-thirds of it--parts with its oxygen in
order to oxidise the remaining third.[36] From an intermediate substance,
RX, two extremes, R and RX_{3} are formed, just as nitrous anhydride
splits up into nitric oxide and nitric anhydride (or nitric acid). The
resulting salt, MClO_{3}, corresponds with _chloric acid_ and potassium
chlorate, KClO_{3}. It is evident that a similar salt may be obtained
directly by the action of chlorine on an alkali if its solution be
heated, because RClO will be first formed, and then RClO_{3}; for
example, 6KHO + 3Cl_{2} = KClO_{3} + 5KCl + 3H_{2}O. Chlorates are so
prepared; for instance, _potassium chlorate_, which is easily separated
from potassium chloride, being sparingly soluble in cold water.[37]
[36] Hydrochloric acid, which is an example of compounds of this kind,
is a saturated substance which does not combine directly with
oxygen, but in which, nevertheless, a considerable quantity of
oxygen may be inserted between the elements forming it. The same
may be observed in a number of other cases. Thus oxygen may be
added or inserted between the elements, sometimes in considerable
quantities, in the saturated hydrocarbons; for instance, in
C_{3}H_{8}, three atoms of oxygen produce an alcohol, glycerin or
glycerol, C_{3}H_{5}(OH)_{3}. We shall meet with similar examples
hereafter. This is generally explained by regarding oxygen as a
bivalent element--that is, as capable of combining with two
different elements, such as chlorine, hydrogen, &c. On the basis
of this view, it may be inserted between each pair of combined
elements; the oxygen will then be combined with one of the
elements by one of its affinities and with the other element by
its other affinity. This view does not, however, express the
entire truth of the matter, even when applied to the compounds of
chlorine. Hypochlorous acid, HOCl--that is, hydrochloric acid in
which one atom of oxygen is inserted--is, as we have already seen,
a substance of small stability; it might therefore be expected
that on the addition of a fresh quantity of oxygen, a still less
stable substance would be obtained, because, according to the
above view, the chlorine and hydrogen, which form such a stable
compound together, are then still further removed from each other.
But it appears that chloric and perchloric acid, HClO_{3} and
HClO_{4}, are much more stable substances. Furthermore, the
addition of oxygen has also its limit, it can only be added to a
certain extent. If the above representation were true and not
merely hypothetical, there would be no limit to the combination of
oxygen, and the more it entered into one continuous chain the more
unstable would be the resultant compound. But not more than four
atoms of oxygen can be added to hydrogen sulphide, nor to
hydrochloric acid, nor to hydrogen phosphide. This peculiarity
must lie in the properties of oxygen itself; four atoms of oxygen
seem to have the power of forming a kind of radicle which retains
two or several atoms of various other substances--for example,
chlorine and hydrogen, hydrogen and sulphur, sodium and manganese,
phosphorus and metals, &c., forming comparatively stable
compounds, NaClO_{4}, Na_{2}SO_{4}, NaMnO_{4}, Na_{3}PO_{4}, &c.
_See_ Chapter X. Note 1 and Chapter XV.
[37] If chlorine be passed through a _cold_ solution of potash, a
bleaching compound, potassium chloride and hypochlorite, KCl +
KClO, is formed, but if it be passed through a _hot_ solution
potassium chlorate is formed. As this is sparingly soluble in
water, it chokes the gas-conducting tube, which should therefore
be widened out at the end.
Potassium chlorate is usually obtained on a large scale from
calcium chlorate, which is prepared by passing chlorine (as long
as it is absorbed) into water containing lime, the mixture being
kept warm. A mixture of calcium chlorate and chloride is thus
formed in the solution. Potassium chloride is then added to the
warm solution, and on cooling a precipitate of potassium chlorate
is formed as a substance which is sparingly soluble in cold water,
especially in the presence of other salts. The double
decomposition taking place is Ca(ClO_{3})_{2} + 2KCl = CaCl_{2} +
2KClO_{3}. On a small scale in the laboratory potassium chlorate
is best prepared from a strong solution of bleaching powder by
passing chlorine through it and then adding potassium chloride.
KClO_{3} is always formed by the action of an electric current on
a solution of KCl, especially at 80° (Häussermann and Naschold,
1894), so that this method is now used on a large scale.
Potassium chlorate crystallises easily in large colourless tabular
crystals. Its solubility in 100 parts of water at 0° = 3 parts,
20° = 8 parts, 40° = 14 parts, 60° = 25 parts, 80° = 40 parts. For
comparison we will cite the following figures showing the
solubility of potassium chloride and perchlorate in 100 parts of
water: potassium chloride at O° = 28 parts, 20° = 35 parts, 40° =
40 parts, 100° = 57 parts; potassium perchlorate at 0° about 1
part, 20° about 1-3/4 part, 100° about 18 parts. When heated,
potassium chlorate melts (the melting point has been given as from
335°-376°; according to the latest determination by Carnelley,
359°) and decomposes with the evolution of oxygen, potassium
perchlorate being at first formed, as will afterwards be described
(_see_ Note 47). A mixture of potassium chlorate and nitric and
hydrochloric acids effects oxidation and chlorination in
solutions. It deflagrates when thrown upon incandescent carbon,
and when mixed with sulphur (1/3 by weight) it ignites it on being
struck, in which case an explosion takes place. The same occurs
with many metallic sulphides and organic substances. Such mixtures
are also ignited by a drop of sulphuric acid. All these effects
are due to the large amount of oxygen contained in potassium
chlorate, and to the ease with which it is evolved. A mixture of
two parts of potassium chlorate, one part of sugar, and one part
of yellow prussiate of potash acts like gunpowder, but burns too
rapidly, and therefore bursts the guns, and it also has a very
strong oxidising action on their metal. The sodium salt,
NaClO_{3}, is much more soluble than the potassium salt, and it is
therefore more difficult to free it from sodium chloride, &c. The
barium salt is also more soluble than the potassium salt; O° = 24
parts, 20° = 37 parts, 80° = 98 parts of salt per 100 of water.
If dilute sulphuric acid be added to a solution of potassium chlorate,
_chloric acid_ is liberated, but it cannot be separated by distillation,
as it is decomposed in the process. To obtain the free acid, sulphuric
acid must be added to a solution of barium chlorate.[38] The sulphuric
acid gives a precipitate of barium sulphate, and free chloric acid
remains in solution. The solution may be evaporated under the receiver of
an air-pump. This solution is colourless, has no smell, and acts as a
powerful acid (it neutralises sodium hydroxide, decomposes sodium
carbonate, gives hydrogen with zinc, &c.); when heated above 40°,
however, it decomposes, forming chlorine, oxygen, and perchloric acid:
4HClO_{3} = 2HClO_{4} + H_{2}O + Cl_{2} + O_{3}. In a concentrated
condition the acid acts as an exceedingly energetic oxidiser, so that
organic substances brought into contact with it burst into flame. Iodine,
sulphurous acid, and similar oxidisable substances form higher oxidation
products and reduce the chloric acid to hydrochloric acid. Hydrochloric
acid gas gives chlorine with chloric acid (and consequently with
KClO_{3} also) acting in the same manner as it acts on the lower acids:
HClO_{3} + 5HCl = 3H_{2}O + 3Cl_{2}.
[38] Barium chlorate, Ba(ClO_{3})_{2},H_{2}O, is prepared in the
following way: impure chloric acid is first prepared and saturated
with baryta, and the barium salt purified by crystallisation. The
impure free chloric acid is obtained by converting the potassium
in potassium chlorate into an insoluble salt. This is done by
adding tartaric or hydrofluosilicic acid to a solution of
potassium chlorate, because potassium tartrate and potassium
silicofluoride are very sparingly soluble in water. Chloric acid
is easily soluble in water.
By cautiously acting on potassium chlorate with sulphuric acid, the
_dioxide_ (_chloric peroxide_), ClO_{2},[39] is obtained (Davy, Millon).
This gas is easily liquefied in a freezing mixture, and boils at +10°.
The vapour density (about 35 if H = 1) shows that the molecule of this
substance is ClO_{2}.[40] In a gaseous or liquid state it very easily
explodes (for instance, at 60°, or by contact with organic compounds or
finely divided substances, &c.), forming Cl and O_{2}, and in many
instances[41] therefore it acts as an oxidising agent, although (like
nitric peroxide) it may itself be further oxidised.[42] In dissolving in
water or alkalis chloric peroxide gives chlorous and hypochlorous
acids--2ClO_{2} + 2KHO = KClO_{3} + KClO_{2} + H_{2}O--and therefore,
like nitric peroxide, the dioxide may be regarded as an intermediate
oxide between the (unknown) anhydrides of chlorous and chloric acids:
4ClO_{2} = Cl_{2}O_{3} + Cl_{2}O_{3}.[43]
[39] To prepare ClO_{2} 100 grams of sulphuric acid are cooled in a
mixture of ice and salt, and 15 grams of powdered potassium
chlorate are gradually added to the acid, which is then carefully
distilled at 20° to 40°, the vapour given off being condensed in a
freezing mixture. Potassium perchlorate is then formed: 3KClO_{3}
+ 2H_{2}SO_{4} = 2KHSO_{4} + KClO_{4} + 2ClO_{2} + H_{2}O. The
reaction may result in an explosion. Calvert and Davies obtained
chloric peroxide without the least danger by heating a mixture of
oxalic acid and potassium chlorate in a test tube in a water-bath.
In this case 2KClO_{3} + 3C_{2}H_{2}O_{4},2H_{2}O = 2C_{2}HKO_{4}
+ 2CO_{2} + 2ClO_{2} + 8H_{2}O. The reaction is still further
facilitated by the addition of a small quantity of sulphuric acid.
If a solution of HCl acts upon KClO_{3} at the ordinary
temperature, a mixture of Cl_{2} and ClO_{2} is formed, but if the
temperature be raised to 80° the greater part of the ClO_{2}
decomposes, and when passed through a hot solution of MnCl_{2} it
oxidises it. Gooch and Kreider proposed (1894) to employ this
method for preparing small quantities of chlorine in the
laboratory.
[40] By analogy with nitric peroxide it might be expected that at low
temperatures a doubling of the molecule into Cl_{2}O_{4} would
take place, as the reactions of ClO_{2} point to its being a mixed
anhydride of HClO_{2} and HClO_{3}.
[41] Owing to the formation of this chlorine dioxide, a mixture of
potassium chlorate and sugar is ignited by a drop of sulphuric
acid. This property was formerly made use of for making matches,
and is now sometimes employed for setting fire to explosive
charges by means of an arrangement in which the acid is caused to
fall on the mixture at the moment required. An interesting
experiment on the combustion of phosphorus under water may be
conducted with chlorine dioxide. Pieces of phosphorus and of
potassium chlorate are placed under water, and sulphuric acid is
poured on to them (through a long funnel); the phosphorus then
burns at the expense of the chlorine dioxide.
[42] Potassium permanganate oxidises chlorine dioxide into chloric acid
(Fürst).
[43] The euchlorine obtained by Davy by gently heating potassium
chlorate with hydrochloric acid is (Pebal) a mixture of chlorine
dioxide and free chlorine. The liquid and gaseous chlorine oxide
(Note 35), which Millon considered to be Cl_{2}O_{3}, probably
contains a mixture of ClO_{2} (vapour density 35), Cl_{2}O_{3}
(whose vapour density should be 59), and chlorine (vapour density
35·5), since its vapour density was determined to be about 40.
As the salts of chloric acid, HClO_{3}, are produced by the splitting
up of the salts of hypochlorous acid, so in the same way the salts
of perchloric acid, HClO_{4}, are produced from the salts of chloric
acid, HClO_{3}. But this is the highest form of the oxidation of HCl.
_Perchloric acid_, HClO_{4}, is the most stable of all the acids of
chlorine. When fused potassium chlorate begins to swell up and solidify,
after having parted with one-third of its oxygen, potassium chloride and
potassium perchlorate have been formed according to the equation 2KClO_{3}
= KClO_{4} + KCl + O_{2}.
The formation of this salt is easily observed in the preparation of
oxygen from potassium chlorate, owing to the fact that the potassium
perchlorate fuses with greater difficulty than the chlorate, and
therefore appears in the molten salt as solid grains (_see_ Chapter III.
Note 12). Under the action of certain acids--for instance, sulphuric and
nitric--potassium chlorate also gives potassium perchlorate. This latter
may be easily purified, because it is but sparingly soluble in water,
although all the other salts of perchloric acid are very soluble and even
deliquesce in the air. The perchlorates, although they contain more
oxygen than the chlorates, are decomposed with greater difficulty, and
even when thrown on ignited charcoal give a much feebler deflagration
than the chlorates. Sulphuric acid (at a temperature not below 100°)
evolves volatile and to a certain extent stable perchloric acid from
potassium perchlorate. Neither sulphuric nor any other acid will further
decompose perchloric acid as it decomposes chloric acid. Of all the acids
of chlorine, perchloric acid alone can be distilled.[44] The pure hydrate
HClO_{4}[45] is a colourless and exceedingly caustic substance which
fumes in the air and has a specific gravity 1·78 at 15° (sometimes, after
being kept for some time, it decomposes with a violent explosion). It
explodes violently when brought into contact with charcoal, paper, wood,
and other organic substances. If a small quantity of water be added to
this hydrate, and it be cooled, a crystallo-hydrate, ClHO_{4},H_{2}O,
separates out. This is much more stable, but the liquid hydrate
HClO_{4},2H_{2}O is still more so. The acid dissolves in water in all
proportions, and its solutions are distinguished for their stability.[46]
When ignited both the acid and its salts are decomposed, with the
evolution of oxygen.[47]
[44] If a solution of chloric acid, HClO_{3}, be first concentrated
over sulphuric acid under the receiver of an air-pump and
afterwards distilled, chlorine and oxygen are evolved and
perchloric acid is formed: 4HClO_{3} = 2HClO_{4} + Cl_{2} + 3O +
H_{2}O. Roscoe accordingly decomposed directly a solution of
potassium chlorate by hydrofluosilicic acid, decanted it from the
precipitate of potassium silicofluoride, K_{2}SiF_{6},
concentrated the solution of chloric acid, and then distilled it,
perchloric acid being then obtained (_see_ following footnote).
That chloric acid is capable of passing into perchloric acid is
also seen from the fact that potassium permanganate is
decolorised, although slowly, by the action of a solution of
chloric acid. On decomposing a solution of potassium chlorate by
the action of an electric current, potassium perchlorate is
obtained at the positive electrode (where the oxygen is evolved).
Perchloric acid is also formed by the action of an electric
current on solutions of chlorine and chlorine monoxide. Perchloric
acid was obtained by Count Stadion and afterwards by Serullas, and
was studied by Roscoe and others.
[45] Perchloric acid, which is obtained in a free state by the action
of sulphuric acid on its salts, may be separated from a solution
very easily by distillation, being volatile, although it is
partially decomposed by distillation. The solution obtained after
distillation may be concentrated by evaporation in open vessels.
In the distillation the solution reaches a temperature of 200°,
and then a very constant liquid hydrate of the composition
HClO_{4},2H_{2}O is obtained in the distillate. If this hydrate be
mixed with sulphuric acid, it begins to decompose at 100°, but
nevertheless a portion of the acid passes over into the receiver
without decomposing, forming a crystalline hydrate HClO_{4},H_{2}O
which melts at 50°. On carefully heating this hydrate it breaks up
into perchloric acid, which distills over below 100°, and into the
liquid hydrate HClO_{4},2H_{2}O. The acid HClO_{4} may also be
obtained by adding one-fourth part of strong sulphuric acid to
potassium chlorate, carefully distilling and subjecting the
crystals of the hydrate HClO_{4},H_{2}O obtained in the distillate
to a fresh distillation. Perchloric acid, HClO_{4}, itself does
not distil, and is decomposed on distillation until the more
stable hydrate HClO_{4}, H_{2}O is formed; this decomposes into
HClO_{4} and HClO_{4},2H_{2}O, which latter hydrate distils without
decomposition. This forms an excellent example of the influence of
water on stability, and of the property of chlorine of giving
compounds of the type ClX_{7}, of which all the above hydrates,
ClO_{3}(OH), ClO_{2}(OH)_{3}, and ClO(OH)_{5}, are members.
Probably further research will lead to the discovery of a hydrate
Cl(OH)_{7}.
[46] According to Roscoe the specific gravity of perchloric acid = 1·782
and of the hydrate HClO_{4},H_{2}O in a liquid state (50°) 1·811;
hence a considerable contraction takes place in the combination of
HClO_{4} with H_{2}O.
[47] The decomposition of salts analogous to potassium chlorate has been
more fully studied in recent years by Potilitzin and P. Frankland.
Professor Potilitzin, by decomposing, for example, lithium
chlorate LiClO_{3}, found (from the quantity of lithium chloride
and oxygen) that at first the decomposition of the fused salt
(368°) takes place according to the equation, 3LiClO_{3} = 2LiCl +
LiClO_{4} + 5O, and that towards the end the remaining salt is
decomposed thus: 5LiClO_{3} = 4LiCl + LiClO_{4} + 10O. The
phenomena observed by Potilitzin obliged him to admit that lithium
perchlorate is capable of decomposing simultaneously with lithium
chlorate, with the formation of the latter salt and oxygen; and
this was confirmed by direct experiment, which showed that lithium
chlorate is always formed in the decomposition of the perchlorate.
Potilitzin drew particular attention to the fact that the
decomposition of potassium chlorate and of salts analogous to it,
although exothermal (Chapter III., Note 12), not only does not
proceed spontaneously, but requires time and a rise of temperature
in order to attain completion, which again shows that chemical
equilibria are not determined by the heat effects of reactions
only.
P. Frankland and J. Dingwall (1887) showed that at 448° (in the
vapour of sulphur) a mixture of potassium chlorate and powdered
glass is decomposed almost in accordance with the equation
2KClO_{3} = KClO_{4} + KCl + O_{2}, whilst the salt by itself
evolves about half as much oxygen, in accordance with the
equation, 8KClO_{3} = 5KClO_{4} + 3KCl + 2O_{2}. The decomposition
of potassium perchlorate in admixture with manganese peroxide
proceeds to completion, KClO_{4} = KCl + 2O_{2}. But in
decomposing by itself the salt at first gives potassium chlorate,
approximately according to the equation 7KClO_{4} = 2KClO_{3} +
5KCl + 11O_{2}. Thus there is now no doubt that when potassium
chlorate is heated, the perchlorate is formed, and that this salt,
in decomposing with evolution of oxygen, again gives the former
salt.
In the decomposition of barium hypochlorite, 50 per cent. of the
whole amount passes into chlorate, in the decomposition of
strontium hypochlorite (Potilitzin, 1890) 12·5 per cent., and of
calcium hypochlorite about 2·5 per cent. Besides which Potilitzin
showed that the decomposition of the hypochlorites and also of the
chlorates is always accompanied by the formation of a certain
quantity of the oxides and by the evolution of chlorine, the
chlorine being displaced by the oxygen disengaged. Spring and
Prost (1889) represent the evolution of oxygen from KClO_{3} as
due to the salt first splitting up into base and anhydride, thus
(1) 2MClO_{3} = M_{2}O + Cl_{2}O_{5}; (2) Cl_{2}O_{5} = Cl_{2} +
O_{3}; and (3) M_{2}O + Cl = 2MCl + O.
I may further remark that the decomposition of potassium chlorate
as a reaction evolving heat easily lends itself for this very
reason to the contact action of manganese peroxide and other
similar admixtures; for such very feeble influences as those of
contact may become evident either in those cases (for instance,
detonating gas, hydrogen peroxide, &c.), when the reaction is
accompanied by the evolution of heat, or when (for instance, H_{2}
+ I_{2}, &c.) little heat is absorbed or evolved. In these cases
it is evident that the existing equilibrium is not very stable,
and that a small alteration in the conditions at the surfaces of
contact may suffice to upset it. In order to conceive the _modus
operandi_ of contact phenomena, it is enough to imagine, for
instance, that at the surface of contact the movement of the atoms
in the molecules changes from a circular to an elliptical path.
Momentary and transitory compounds may he formed, but their
formation cannot affect the explanation of the phenomena.
On comparing chlorine as an element not only with nitrogen and carbon but
with all the other non-metallic elements (chlorine has so little analogy
with the metals that a comparison with them would be superfluous), we
find in it the following fundamental properties of _the halogens_ or
salt-producers. With metals chlorine gives salts (such as sodium
chloride, &c.); with hydrogen a very energetic and monobasic acid HCl,
and the same quantity of chlorine is able by metalepsis to replace the
hydrogen; with oxygen it forms unstable oxides of an acid character.
These properties of chlorine are possessed by three other elements,
bromine, iodine, and fluorine. They are members of one natural family.
Each representative has its peculiarities, its individual properties and
points of distinction, in combination and in the free state--otherwise
they would not be independent elements; but the repetition in all of them
of the same chief characteristics of the family enables one more quickly
to grasp all their various properties and to classify the elements
themselves.
In order to have a guiding thread in forming comparisons between the
elements, attention must however be turned not only to their points of
resemblance but also to those of their properties and characters in
which they differ most from each other. And the atomic weights of the
elements must be considered as their most elementary property, since
this is a quantity which is most firmly established, and must be taken
account of in all the reactions of the element. The halogens have the
following atomic weights--
F = 19, Cl = 35·5, Br = 80, I = 127.
All the properties, physical and chemical, of the elements and their
corresponding compounds must evidently be in a certain dependence on
this fundamental point, if the grouping in one family be natural.[47 bis]
And we find in reality that, for instance, the properties of bromine,
whose atomic weight is almost the mean between those of iodine and
chlorine, occupy a mean position between those of these two elements. The
second measurable property of the elements is their equivalence or their
capacity for forming _compounds of definite forms_. Thus carbon or
nitrogen in this respect differs widely from the halogens. Although the
form ClO_{2} corresponds with NO_{2} and CO_{2}, yet the last is the
highest oxide of carbon, whilst that of nitrogen is N_{2}O_{5}, and for
chlorine, if there were an anhydride of perchloric acid, its composition
would be Cl_{2}O_{7}, which is quite different from that of carbon. In
respect to the forms of their compounds the halogens, like all elements
of one family or group, are perfectly analogous to each other, as is seen
from their hydrogen compounds:
HF, HCl, HBr, HI.
[47 bis] See, for example the melting point of NaCl, NaBr, NaI in
Chapter II. Note 27. According to F. Freyer and V. Meyer (1892),
the following are the boiling points of some of the corresponding
compounds of chlorine and bromine:
BCl_{5} 17° BBr_{3} 90°
SiCl_{3} 59° SiBr_{4} 153°
PCl_{3} 76° PBr_{3} 175°
SbCl_{3} 223° SbBr_{3} 275°
BiCl_{3} 447° BiBr_{3} 453°
SnCl_{4} 606° SnBr_{4} 619°
ZnCl_{2} 730° ZnBr_{2} 650°
Thus for all the more volatile compounds the replacement of
chlorine by bromine raises the boiling point, but in the ease of
ZnX_{2} it lowers it (Chapter XV. Note 19).
Their oxygen compounds exhibit a similar analogy. Only fluorine does not
give any oxygen compounds. The iodine and bromine compounds corresponding
with HClO_{3} and HClO_{4} are HBrO_{3} and HBrO_{4}, HIO_{3} and
HIO_{4}. On comparing the properties of these acids we can even predict
that fluorine will not form any oxygen compound. For iodine is easily
oxidised--for instance, by nitric acid--whilst chlorine is not directly
oxidised. The oxygen acids of iodine are comparatively more stable than
those of chlorine; and, generally speaking, the affinity of iodine for
oxygen is much greater than that of chlorine. Here also bromine occupies
an intermediate position. In fluorine we may therefore expect a still
smaller affinity for oxygen than in chlorine--and up to now it has not
been combined with oxygen. If any oxygen compounds of fluorine should be
obtained, they will naturally be exceedingly unstable. The relation of
these elements to hydrogen is the reverse of the above. Fluorine has so
great an affinity for hydrogen that it decomposes water at the ordinary
temperature; whilst iodine has so little affinity for hydrogen that
hydriodic acid, HI, is formed with difficulty, is easily decomposed, and
acts as a reducing agent in a number of cases.
From the form of their compounds the halogens are _univalent elements_
with respect to hydrogen and septivalent with respect to oxygen, N being
trivalent to hydrogen (it gives NH_{3}) and quinqui-valent to oxygen (it
gives N_{2}O_{5}), and C being quadrivalent to both H and O as it forms
CH_{4} and CO_{2}. And as not only their oxygen compounds, but also their
hydrogen compounds, have acid properties, the halogens are _elements_
of an exclusively _acid character_. Such metals as sodium, potassium,
barium only give basic oxides. In the case of nitrogen, although it
forms acid oxides, still in ammonia we find that capacity to give an
alkali with hydrogen which indicates a less distinctly acid character
than in the halogens. In no other elements is the acid-giving property
so strongly developed as in the halogens.
In describing certain peculiarities characterising the halogens, we
shall at every step encounter a confirmation of the above-mentioned
general relations.
As _fluorine_ decomposes water with the evolution of oxygen, F_{2} +
H_{2}O = 2HF + O, for a long time all efforts to obtain it in free state
by means of methods similar to those for the preparation of chlorine
proved fruitless.[48] Thus by the action of hydrofluoric acid on
manganese peroxide, or by decomposing a solution of hydrofluoric acid by
an electric current, either oxygen or a mixture of oxygen and fluorine
were obtained instead of fluorine. Probably a certain quantity of
fluorine[48 bis] was set free by the action of oxygen or an electric
current on incandescent and fused calcium fluoride, but at a high
temperature fluorine acts even on platinum, and therefore it was not
obtained. When chlorine acted on silver fluoride, AgF, in a vessel of
natural fluor spar, CaF_{2}, fluorine was also liberated; but it was
mixed with chlorine, and it was impossible to study the properties of
the resultant gas. Brauner (1881) also obtained fluorine by igniting
cerium fluoride, 2CeF_{4} = 2CeF_{3} + F_{2}; but this, like all
preceding efforts, only showed fluorine to be a gas which decomposes
water, and is capable of acting in a number of instances like chlorine,
but gave no possibility of testing its properties. It was evident that it
was necessary to avoid as far as possible the presence of water and a
rise of temperature; this Moissan succeeded in doing in 1886. He
decomposed anhydrous hydrofluoric acid, liquefied at a temperature of
-23° and contained in a U-shaped tube (to which a small quantity of
potassium fluoride had been added to make it a better conductor), by the
action of a powerful electric current (twenty Bunsen's elements in
series). Hydrogen was then evolved at the negative pole, and fluorine
appeared at the positive pole (of iridium platinum) as a pale green gas
which decomposed water with the formation of ozone and hydrofluoric acid,
and combined directly with silicon (forming silicon fluoride, SiF_{4}),
boron (forming BF_{3}), sulphur, &c. Its density (H = 1) is 18, so that
its molecule is F_{2}. But the action of fluorine on metals at the
ordinary temperature is comparatively feeble, because the metallic
fluoride formed coats the remaining mass of the metals; it is, however,
completely absorbed by iron. Hydrocarbons (such as naphtha), alcohol,
&c., immediately absorb fluorine, with the formation of hydrofluoric
acid. Fluorine when mixed with hydrogen can easily be made to explode
violently, forming hydrofluoric acid.[49]
[48] Even before free fluorine was obtained (1886) it was evident from
experience gained in the efforts made to obtain it, and from
analogy, that it would decompose water (_see_ first Russian
edition of the _Principles of Chemistry_).
[48 bis] It is most likely that in this experiment of Fremy's, which
corresponds with the action of oxygen on calcium chloride,
fluorine was set free, but that a converse reaction also
proceeded, CaO + F_{2} = CaF_{2} + O--that is, the calcium
distributed itself between the oxygen and fluorine. MnF_{4}, which
is capable of splitting up into MnF_{2} and F_{2}, is without
doubt formed by the action of a strong solution of hydrofluoric
acid on manganese peroxide, but under the action of water the
fluorine gives hydrofluoric acid, and probably this is aided by
the affinity of the manganese fluoride and hydrofluoric acid. In
all the attempts made (by Davy, Knox, Louget, Fremy, Gore, and
others) to decompose fluorides (those of lead, silver, calcium,
and others) by chlorine, there were doubtless also cases of
distribution, a portion of the metal combined with chlorine and a
portion of the fluorine was evolved; but it is improbable that any
decisive results were obtained. Fremy probably obtained fluorine,
but not in a pure state.
[49] According to Moissan, fluorine is disengaged by the action of an
electric current on fused hydrogen potassium fluoride, KHF_{2}.
The present state of chemical knowledge is such that the knowledge
of the properties of an element is much more general than the
knowledge of the free element itself. It is useful and
satisfactory to learn that even fluorine in the free state has not
succeeded in eluding experiment and research, that the efforts to
isolate it have been crowned with success, but the sum total of
chemical data concerning fluorine as an element gains but little
by this achievement. The gain will, however, be augmented if it be
now possible to subject fluorine to a comparative study in
relation to oxygen and chlorine. There is particular interest in
the phenomena of the distribution of fluorine and oxygen, or
fluorine and chlorine, competing under different conditions and
relations. We may add that Moissan (1892) found that free fluorine
decomposes H_{2}S, HCl, HBr, CS_{2}, and CNH with a flash; it does
not act upon O_{2}, N_{2}, CO, and CO_{2}; Mg, Al, Ag, and Ni,
when heated, burn in it, as also do S, Se, P (forms PF_{5}); it
reacts upon H_{2} even in the dark, with the evolution of 366·00
units of heat. At a temperature of -95°, F_{2} still retains its
gaseous state. Soot and carbon in general (but not the diamond)
when heated in gaseous fluorine form _fluoride of carbon_, CF_{4}
(Moissan, 1890); this compound is also formed at 300° by the
double decomposition of CCl_{4} and AgF; it is a gas which
liquefies at 10° under a pressure of 5 atmospheres. With an
alcoholic solution of KHO, CF_{4} gives K_{2}CO_{3}, according to
the equation CF_{4} + 6KHO = K_{2}CO_{3} + 4KF + 3H_{2}O. CF_{4}
is not soluble in water, but it is easily soluble in CCl_{4} and
alcohol.
In 1894 Brauner obtained fluorine directly by igniting the easily
formed[49 bis] double lead salt HF,3KF,PbF_{4}, which first, at 230°,
decomposes with the evolution of HF, and then splits up forming
3KF,PbF_{2} and fluorine F_{2}, which is recognised by the fact that it
liberates iodine from KI and easily combines with silicon, forming
SiF_{4}. This method gives chemically pure fluorine, and is based upon
the breaking up of the higher compound--tetrafluoride of lead, PbF_{4},
corresponding to PbO_{2}, into free fluorine, F_{2}, and the lower more
stable form--bifluoride of lead, PbF_{2}, which corresponds to PbO; that
is, this method resembles the ordinary method of obtaining chlorine by
means of MnO_{2}, as MnCl_{4} here breaks up into MnCl_{2} and chlorine,
just as PbF_{4} splits up into PbF_{2} and fluorine.
[49 bis] T. Nikolukin (1885) and subsequently Friedrich and Classen
obtained PbCl_{4} and a double ammonium salt of tetrachloride of
lead (starting from the binoxide), PbCl_{4}2NH_{4}Cl; Hutchinson
and Pallard obtained a similar salt of acetic acid (1893)
corresponding to PbX_{4} by treating red lead with strong acetic
acid; the composition of this salt is Pb(C_{2}H_{3}O_{2})_{4}; it
melts (and decomposes) at about 175°. Brauner (1894) obtained a
salt corresponding to tetrafluoride of lead, PbF_{4}, and the acid
corresponding to it, H_{4}PbF_{8}. For example, by treating
potassium plumbate (Chapter XVIII. Note 55) with strong HF, and
also the above-mentioned tetra-acetate with a solution of KHF_{2},
Brauner obtained crystalline HK_{3}PbF_{8}--i.e. the salt from
which he obtained fluorine.
Among the compounds of fluorine, calcium fluoride, CaF_{2}, is somewhat
widely distributed in nature as fluor spar,[50] whilst _cryolite_, or
aluminium sodium fluoride, Na_{3}AlF_{6}, is found more rarely (in large
masses in Greenland). Cryolite, like fluor spar, is also insoluble in
water, and gives hydrofluoric acid with sulphuric acid. Small quantities
of fluorine have also in a number of cases been found in the bodies of
animals, in the blood, urine, and bones. If fluorides occur in the bodies
of animals, they must have been introduced in food, and must occur in
plants and in water. And as a matter of fact river, and especially sea,
water always contains a certain, although small, quantity of fluorine
compounds.
[50] It is called spar because it very frequently occurs as crystals of
a clearly laminar structure, and is therefore easily split up into
pieces bounded by planes. It is called fluor spar because when
used as a flux it renders ores fusible, owing to its reacting with
silica, SiO_{2} + 2CaF_{2} = 2CaO + SiF_{4}; the silicon fluoride
escapes as a gas and the lime combines with a further quantity of
silica, and gives a vitreous slag. Fluor spar occurs in mineral
veins and rocks, sometimes in considerable quantities. It always
crystallises in the cubic system, sometimes in very large
semi-transparent cubic crystals, which are colourless or of
different colours. It is insoluble in water. It melts under the
action of heat, and crystallises on cooling. The specific gravity
is 3·1. When steam is passed over incandescent fluor spar, lime
and hydrofluoric acid are formed: CaF_{2} + H_{2}O = CaO + 2HF. A
double decomposition is also easily produced by fusing fluor spar
with sodium or potassium hydroxides, or potash, or even with their
carbonates; the fluorine then passes over to the potassium or
sodium, and the oxygen to the calcium. In solutions--for example,
Ca(NO_{3})_{2} + 2KF = CaF_{2} (precipitate) + 2KNO_{3} (in
solution)--the formation of calcium fluoride takes place, owing to
its very sparing solubility. 26,000 parts of water dissolve one
part of fluor spar.
Hydrifluoric acid, HF, cannot be obtained from fluor spar in glass
retorts, because glass is acted on by and destroys the acid. It is
prepared in lead vessels, and when it is required pure, in platinum
vessels, because lead also acts on hydrofluoric acid, although only very
feebly on the surface, and when once a coating of fluoride and sulphate
of lead is formed no further action takes place. Powdered fluor spar and
sulphuric acid evolve hydrofluoric acid (which fumes in the air) even at
the ordinary temperature, CaF_{2} + H_{2}SO_{4} = CaSO_{4} + 2HF. At 130°
fluor spar is completely decomposed by sulphuric acid. The acid is then
evolved as vapour, which may be condensed by a freezing mixture into an
anhydrous acid. The condensation is aided by pouring water into the
receiver of the condenser, as the acid is easily soluble in cold water.
In the liquid anhydrous form hydrofluoric acid boils at +19°, and its
specific gravity at 12·8° = 0·9849.[51] It dissolves in water with the
evolution of a considerable amount of heat, and gives a solution of
constant boiling point which distils over at 120°; showing that the acid
is able to combine with water. The specific gravity of the compound is
1·15, and its composition HF,2H_{2}O.[52] With an excess of water a
dilute solution distils over first. The aqueous solution and the acid
itself must be kept in platinum vessels, but the dilute acid may be
conveniently preserved in vessels made of various organic materials,
such as gutta-percha, or even in glass vessels having an interior
coating of paraffin. Hydrofluoric acid does not act on hydrocarbons
and many other substances, but it acts in a highly corrosive manner
on metals, glass, porcelain, and the majority of rock substances.[53]
It also attacks the skin, and is distinguished by its poisonous
properties, so that in working with the acid a strong draught must be
kept up, to prevent the possibility of the fumes being inhaled. The
non-metals do not act on hydrofluoric acid, but all metals--with the
exception of mercury, silver, gold, and platinum, and, to a certain
degree, lead--decompose it with the evolution of hydrogen. With bases
it gives directly metallic fluorides, and behaves in many respects like
hydrochloric acid. There are, however, several distinct individual
differences, which are furthermore much greater than those between
hydrochloric, hydrobromic, and hydriodic acids. Thus the silver
compounds of the latter are insoluble in water, whilst silver fluoride
is soluble. Calcium fluoride, on the contrary, is insoluble in water,
whilst calcium chloride, bromide, and iodide are not only soluble, but
attract water with great energy. Neither hydrochloric, hydrobromic, nor
hydriodic acid acts on sand and glass, whilst hydrofluoric acid corrodes
them, forming gaseous silicon fluoride. The other halogen acids only form
normal salts, KCl, NaCl, with Na or K, whilst hydrofluoric acid gives
acid salts, for instance HKF_{2} (and by dissolving KF in liquid HF,
KHF_{2}2HF is obtained). This latter property is in close connection with
the fact that at the ordinary temperature the vapour density of
hydrofluoric acid is nearly 20, which corresponds with a formula
H_{2}F_{2}, as Mallet (1881) showed; but a depolymerisation occurs with a
rise of temperature, and the density approaches 10, which answers to the
formula HF.[54]
[51] According to Gore. Fremy obtained anhydrous hydrofluoric acid by
decomposing lead fluoride at a red heat, by hydrogen, or by
beating the double salt HKF_{2}, which easily crystallises (in
cubes) from a solution of hydrofluoric acid, half of which has
been saturated with potassium hydroxide. Its vapour density
corresponds to the formula HF.
[52] This composition corresponds to the crystallo-hydrate HCl,2H_{2}O.
All the properties of hydrofluoric acid recall those of
hydrochloric acid, and therefore the comparative ease with which
hydrofluoric acid is liquefied (it boils at +19°, hydrochloric
acid at -35°) must be explained by a polymerisation taking place at
low temperatures, as will be afterwards explained, H_{2}F_{2} being
formed, and therefore in a liquid state it differs from
hydrochloric acid, in which a phenomenon of a similar kind has not
yet been observed.
[53] The corrosive action of hydrofluoric acid on glass and similar
siliceous compounds is based upon the fact that it acts on silica,
SiO_{2}, as we shall consider more fully in describing that
compound, forming gaseous silicon fluoride, SiO_{2} + 4HF =
SiF_{4} + 2H_{2}O. Silica, on the other hand, forms the binding
(acid) element of glass and of the mass of mineral substances
forming the salts of silica. When it is removed the cohesion is
destroyed. This is made use of in the arts, and in the laboratory,
for etching designs and scales, &c., on glass. In _engraving on
glass_ the surface is covered with a varnish composed of four
parts of wax and one part of turpentine. This varnish is not acted
on by hydrofluoric acid, and it is soft enough to allow of designs
being drawn upon it whose lines lay bare the glass. The drawing is
made with a steel point, and the glass is afterwards laid in a
lead trough in which a mixture of fluor spar and sulphuric acid is
placed. The sulphuric acid must be used in considerable excess, as
otherwise transparent lines are obtained (owing to the formation
of hydrofluosilicic acid). After being exposed for some time, the
varnish is removed (melted) and the design drawn by the steel
point is found reproduced in dull lines. The drawing may be also
made by the direct application of a mixture of a silicofluoride
and sulphuric acid, which forms hydrofluoric acid.
[54] Mallet (1881) determined the density at 30° and 100°, previous to
which Gore (1869) had determined the vapour density at 100°,
whilst Thorpe and Hambly (1888) made fourteen determinations
between 26° and 88°, and showed that within this limit of
temperature the density gradually diminishes, just like the vapour
of acetic acid, nitrogen dioxide, and others. The tendency of HF
to polymerise into H_{2}F_{2} is probably connected with the
property of many fluorides of forming acid salts--for example,
KHF_{2} and H_{2}SiF_{6}. We saw above that HCl has the same
property (forming, for instance, H_{2}PtCl_{6}, &c., p. 457), and
hence this property of hydrofluoric acid does not stand isolated
from the properties of the other halogens.
The analogy between chlorine and the other two halogens, bromine and
iodine, is much more perfect. Not only have their hydrates or halogen
acids much in common, but they themselves resemble chlorine in many
respects,[55] and even the properties of the corresponding metallic
compounds of bromine and iodine are very much alike. Thus, the chlorides,
bromides, and iodides of sodium and potassium crystallise in the cubic
system, and are soluble in water; the chlorides of calcium, aluminium,
magnesium, and barium are just as soluble in water as the bromides and
iodides of these metals. The iodides and bromides of silver and lead are
sparingly soluble in water, like the chlorides of these metals. The
oxygen compounds of bromine and iodine also present a very strong analogy
to the corresponding compounds of chlorine. A hypobromous acid is known
corresponding with hypochlorous acid. The salts of this acid have the
same bleaching property as the salts of hypochlorous acid. Iodine was
discovered in 1811 by Courtois in kelp, and was shortly afterwards
investigated by Clement, Gay-Lussac, and Davy. Bromine was discovered in
1826 by Balard in the mother liquor of sea water.
[55] For instance, the experiment with Dutch metal foil (Note 16) may
be made with bromine just as well as with chlorine. A very
instructive experiment on the direct combination of the halogens
with metals maybe made by throwing a small piece (a shaving) of
aluminium into a vessel containing liquid bromine; the aluminium,
being lighter, floats on the bromine, and after a certain time
reaction sets in accompanied by the evolution of heat, light, and
fumes of bromine. The incandescent piece of metal moves rapidly
over the surface of the bromine in which the resultant aluminium
bromide dissolves. For the sake of comparison we will proceed to
cite several thermochemical data (Thomsen) for analogous actions
of (1) chlorine, (2) bromine, and (3) iodine, with respect to
metals; the halogen being expressed by the symbol X, and the plus
sign connecting the reacting substances. All the figures are given
in thousands of calories, and refer to molecular quantities in
grams and to the ordinary temperature:--
1 2 3
K_{2} + X_{2} 211 191 160
Na_{2} + X_{2} 195 172 138
Ag_{2} + X_{2} 59 45 28
Hg_{2} + X_{2} 83 68 48
Hg + X_{2} 63 51 34
Ca + X_{2} 170 141 --
Ba + X_{2} 195 170 --
Zn + X_{2} 97 76 49
Pb + X_{2} 83 64 40
Al + X_{2} 161 120 70
We may remark that the latent heat of vaporisation of the
molecular weight Br_{2} is about 7·2, and of iodine 6·0 thousand
heat units, whilst the latent heat of fusion of Br_{2} is about
0·3, and of I_{2} about 3·0 thousand heat units. From this it is
evident that the difference between the amounts of heat evolved
does not depend on the difference in physical state. For instance,
the vapour of iodine in combining with Zn to form ZnI_{2} would
give 48 + 8 + 3, or about sixty thousand heat units, or 1-1/2
times less than Zn + Cl_{2}.
_Bromine_ and iodine, like chlorine, occur in sea water in combination
with metals. However, the amount of bromides, and especially of iodides,
in sea water is so small that their presence can only be discovered by
means of sensitive reactions.[56] In the extraction of salt from sea
water the bromides remain in the mother liquor. Iodine and bromine also
occur combined with silver, in admixture with silver chloride, as a rare
ore which is mainly found in America. Certain mineral waters (those of
Kreuznach and Staro-rossüsk) contain metallic bromides and iodides,
always in admixture with an excess of sodium chloride. Those upper strata
of the Stassfurt rock salt (Chapter X.) which are a source of potassium
salts also contain metallic bromides,[57] which collect in the mother
liquors left after the crystallisation of the potassium salts; and this
now forms the chief source (together with certain American springs) of
the bromine in common use. Bromine may be easily liberated from a mixture
of bromides and chlorides, owing to the fact that chlorine displaces
bromine from its compounds with sodium, magnesium, calcium, &c. A
colourless solution of bromides and chlorides turns an orange colour
after the passage of chlorine, owing to the disengagement of bromine.[58]
Bromine may be extracted on a large scale by a similar method, but it is
simpler to add a small quantity of manganese peroxide and sulphuric acid
to the mother liquid direct. This sets free a portion of the chlorine,
and this chlorine liberates the bromine.
[56] One litre of sea-water contains about 20 grams of chlorine, and
about 0·07 gram of bromine. The Dead Sea contains about ten times
as much bromine.
[57] But there is no iodine in Stassfurt carnallite.
[58] The chlorine must not, however, be in large excess, as otherwise
the bromine would contain chlorine. Commercial bromine not
unfrequently contains chlorine, as bromine chloride; this is more
soluble in water than bromine, from which it may thus be freed. To
obtain pure bromine the commercial bromine is washed with water,
dried by sulphuric acid, and distilled, the portion coming over at
58° being collected; the greater part is then converted into
potassium bromide and dissolved, and the remainder is added to the
solution in order to separate iodine, which is removed by shaking
with carbon bisulphide. By heating the potassium bromide thus
obtained with manganese peroxide and sulphuric acid, bromine is
obtained quite free from iodine, which, however, is not present in
certain kinds of commercial bromine (the Stassfurt, for instance).
By treatment with potash, the bromine is then converted into a
mixture of potassium bromide and bromate, and the mixture (which
is in the proportion given in the equation) is distilled with
sulphuric acid, bromine being then evolved: 5KBr + KBrO_{3} +
6H_{2}SO_{4} = 6KHSO_{4} + 3H_{2}O + 3Br_{2}. After dissolving the
bromine in a strong solution of calcium bromide and precipitating
with an excess of water, it loses all the chlorine it contained,
because chlorine forms calcium chloride with CaBr_{2}.
Bromine is a _dark brown liquid_, giving brown fumes, and having a
poisonous suffocating smell, whence its name (from the Greek [Greek:
brômos], signifying evil smelling). The vapour density of bromine shows
that its molecule is Br_{2}. In the cold bromine freezes into brown-grey
scales like iodine. The melting point of pure bromine is -7°·05.[59] The
density of liquid bromide at 0° is 3·187, and at 15° about 3·0. The
boiling point of bromine is about 58°·7. Bromine, like chlorine, is
soluble in water; 1 part of bromine at 5° requires 27 parts of water, and
at 15° 29 parts of water. The aqueous solution of bromine is of an
orange colour, and when cooled to -2° yields crystals containing 10
molecules of water to 1 molecule of bromine.[60] Alcohol dissolves a
greater quantity of bromine, and ether a still greater amount. But after
a certain time products of the action of the bromine on these organic
substances are formed in the solutions. Aqueous solutions of the bromides
also absorb a large amount of bromine.
[59] There has long existed a difference of opinion as to the melting
point of pure bromine. By some investigators (Regnault, Pierre) it
was given as between -7° and -8°, and by others (Balard, Liebig,
Quincke, Baumhauer) as between -20° and -25°. There is now no doubt,
thanks more especially to the researches of Ramsay and Young
(1885), that pure bromine melts at about -7°. This figure is not
only established by direct experiment (Van der Plaats confirmed
it), but also by means of the determination of the vapour
tensions. For solid bromine the vapour tension _p_ in mm. at _t_
was found to be--
_p_ = 20 25 30 35 40 45 mm.
_t_ = -16°·6 -14° -12° -10° -8·5° -7°
For liquid bromine--
_p_ = 50 100 200 400 600 760 mm.
_t_ = -5°·0 8°·2 23°·4 40°·4 51°·9 58°·7
These curves intersect at -7°·05. Besides which, in comparing the
vapour tension of many liquids (for example, those given in
Chapter II., Note 27), Ramsay and Young observed that the ratio of
the absolute temperatures (_t_ + 273) corresponding with equal
tension _varies_ for every pair of substances in rectilinear
proportion in dependence upon _t_, and, therefore, for the above
pressure _p_, Ramsay and Young determined the ratio of _t_ + 273
for water and bromine, and found that the straight lines
expressing these ratios for liquid and solid bromine intersect
also at 7°·05; thus, for example, for solid bromine--
_p_ = 20 25 30 35 40 45
273 + _t_ = 256·4 259 261 263 264·6 266
273 + _t_´ = 295·3 299 302·1 304·8 307·2 309·3
_c_ = 1·152 1·154 1·157 1·159 1·161 1·163
where _t_´ indicates the temperature of water corresponding with a
vapour tension _p_, and where _c_ is the ratio of 273 + _t_´ to
273 + _t_. The magnitude of _c_ is evidently expressed with great
accuracy by the straight line _c_ = 1·1703 + 0·0011_t_. In exactly
the same way we find the ratio for liquid bromine and water to be
_c__{1} = 1·1585 + 0·00057t. The intersection of these straight
lines in fact corresponds with -7°·06, which again confirms the
melting point given above for bromine. In this manner it is
possible with the existing store of data to accurately establish
and _verify_ the melting point of substances. Ramsay and Young
established the thermal constants of iodine by exactly the same
method.
[60] The observations made by Paterno and Nasini (by Raoult's method,
Chapter I. Note 49) on the temperature of the formation of ice
(-1°·115, with 1·391 gram of bromine in 100 grams of water) in an
aqueous solution of bromine, showed that bromine is contained in
solutions as the molecule Br_{2}. Similar experiments conducted on
iodine (Kloboukoff 1889 and Beckmann 1890) show that in solution
the molecule is I_{2}.
B. Roozeboom investigated the hydrate of bromine as completely as
the hydrate of chlorine (Notes 9, 10). The temperature of the
complete decomposition of the hydrate is +6°·2; the density of
Br_{2},10H_{2}O = 1·49.
With respect to _iodine_, it is almost exclusively extracted from the
mother liquors after the crystallisation of natural sodium nitrate (Chili
saltpetre) and from the ashes of the sea-weed cast upon the shores of
France, Great Britain, and Spain, sometimes in considerable quantities,
by the high tides. The majority of these sea-weeds are of the genera
_Fucus_, _Laminaria_, &c. The fused ashes of these sea-weeds are called
'kelp' in Scotland and 'varech' in Normandy. A somewhat considerable
quantity of iodine is contained in these sea-weeds. After being burnt
(or subjected to dry distillation) an ash is left which chiefly contains
salts of potassium, sodium, and calcium. The metals occur in the sea-weed
as salts of organic acids. On being burnt these organic salts are
decomposed, forming carbonates of potassium and sodium. Hence, sodium
carbonate is found in the ash of sea plants. The ash is dissolved in hot
water, and on evaporation sodium carbonate and other salts separate, but
a portion of the substances remains in solution. These mother liquors
left after the separation of the sodium carbonate contain chlorine,
bromine, and iodine in combination with metals, the chlorine and iodine
being in excess of the bromine. 13,000 kilos of kelp give about 1,000
kilos of sodium carbonate and 15 kilos of iodine.
The liberation of the iodine from the mother liquor is effected with
comparative ease, because chlorine disengages iodine from potassium
iodide and its other combinations with the metals. Not only chlorine,
but also sulphuric acid, liberates iodine from sodium iodide. Sulphuric
acid, in acting on an iodide, sets hydriodic acid free, but the latter
easily decomposes, especially in the presence of substances capable
of evolving oxygen, such as chromic acid, nitrous acid, and even
ferric salts.[61] Owing to its sparing solubility in water, the iodine
liberated separates as a precipitate. To obtain pure iodine it is
sufficient to distil it, and neglect the first and last portions of
the distillate, the middle portion only being collected. Iodine passes
directly from a state of vapour into a crystalline form, and settles on
the cool portions of the apparatus in tabular crystals, having a black
grey colour and metallic lustre.[62]
[61] In general, 2HI + O = I_{2} + H_{2}O, if the oxygen proceed from a
substance from which it is easily evolved. For this reason
compounds corresponding with the higher stages of oxidation or
chlorination frequently give a lower stage when treated with
hydriodic acid. Ferric oxide, Fe_{2}O_{3}, is a higher oxide, and
ferrous oxide, FeO, a lower oxide; the former corresponds with
FeX_{3}, and the latter with FeX_{2}, and this passage from the
higher to the lower takes place under the action of hydriodic
acid. Thus hydrogen peroxide and ozone (Chapter IV.) are able to
liberate iodine from hydriodic acid. Compounds of copper oxide,
CuO or CuX_{2}, give compounds of the suboxide Cu_{2}O, or CuX.
Even sulphuric acid, which corresponds to the higher stage SO_{3},
is able to act thus, forming the lower oxide SO_{2}. The
liberation of iodine from hydriodic acid proceeds with still
greater ease under the action of substances capable of disengaging
oxygen. In practice, many methods are employed for liberating
iodine from acid liquids containing, for example, sulphuric acid
and hydriodic acid. The higher oxides of nitrogen are most
commonly used; they then pass into nitric oxide. Iodine may even
be disengaged from hydriodic acid by the action of iodic acid, &c.
But there is a limit in these reactions of the oxidation of
hydriodic acid because, under certain conditions, especially in
dilute solutions, the iodine set free is itself able to act as an
oxidising agent--that is, it exhibits the character of chlorine,
and of the halogens in general, to which we shall again have
occasion to refer. In Chili, where a large quantity of iodine is
extracted in the manufacture of Chili nitre, which contains
NaIO_{3}, it is mixed with the acid and normal sulphites of sodium
in solution; the iodine is then precipitated according to the
equation 2NaIO_{3} + 3Na_{2}SO_{3} + 2NaHSO_{3} = 5Na_{2}SO_{4}
+I_{2} + H_{2}O. The iodine thus obtained is purified by
sublimation.
[62] For the final purification of iodine, Stas dissolved it in a
strong solution of potassium iodide, and precipitated it by the
addition of water (_see_ Note 58).
The specific gravity of the crystals of iodine is 4·95. It melts at 114°
and boils at 184°. Its vapour is formed at a much lower temperature, and
is of a violet colour, whence iodine receives its name ([Greek: ioeidês],
violet). The smell of iodine recalls the characteristic smell of
hypochlorous acid; it has a sharp sour taste. It destroys the skin and
organs of the body, and is therefore frequently employed for cauterising
and as an irritant for the skin. In small quantities it turns the skin
brown, but the coloration disappears after a certain time, partly owing
to the volatility of the iodine. Water dissolves only 1/5000 part of
iodine. A brown solution is thus obtained, which bleaches, but much more
feebly than bromine and chlorine. Water which contains salts, and
especially iodides, in solution dissolves iodine in considerable
quantities, and the resultant solution is of a dark brown colour. Pure
alcohol dissolves a small amount of iodine, and in so doing acquires a
brown colour, but the solubility of iodine is considerably increased by
the presence of a small quantity of an iodine compound--for instance,
ethyl iodide--in the alcohol.[63] Ether dissolves a larger amount of
iodine than alcohol; but iodine is particularly soluble in liquid
hydrocarbons, in carbon bisulphide, and in chloroform. A small quantity
of iodine dissolved in carbon bisulphide tints it rose-colour, but in a
somewhat larger amount it gives a violet colour. Chloroform (quite free
from alcohol) is also tinted rose colour by a small amount of iodine.
This gives an easy means for detecting the presence of free iodine in
small quantities. The blue coloration which free iodine gives with
_starch_ may also, as has already been frequently mentioned (_see_
Chapter IV.), serve for the detection of iodine.
[63] The solubility of iodine in solutions containing iodides, and
compounds of iodine in general, may serve, on the one hand, as an
indication that solution is due to a similarity between the
solvent and dissolved substance, and, on the other hand, as an
indirect proof of that view as to solutions which was cited in
Chapter I., because in many instances unstable highly iodised
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.