with oxygen, and we know that ammonia may he oxidised. In these cases
various oxidation products of nitrogen are formed, but in the presence
of water and an excess of oxygen they always give nitric acid. Nitric
acid, as corresponding with the highest oxide, is able, in deoxidising,
to give the lower oxides; it is the only nitrogen acid whose salts occur
somewhat widely in nature, and it has many technical uses, for which
reason we will begin with it.
_Nitric acid_, NHO_{3}, is likewise known as aqua fortis. In a free
state it is only met with in nature in small quantities, in the air
and in rain-water after storms; but even in the atmosphere nitric
acid does not long remain free, but combines with ammonia, traces of
which are always found in air. On falling on the soil and into running
water, &c., the nitric acid everywhere comes into contact with bases
(or their carbonates), which easily act on it, and therefore it is
converted into the nitrates of these bases. Hence nitric acid is always
met with in the form of salts in nature. The soluble salts of nitric
acid are called _nitres_. This name is derived from the Latin _sal
nitri_. The potassium salt, KNO_{3}, is common nitre, and the sodium
salt, NaNO_{3}, Chili saltpetre, or cubic nitre. Nitres are formed
in the soil when a nitrogenous substance is slowly oxidised in the
presence of an alkali by means of the oxygen of the atmosphere. In
nature there are very frequent instances of such oxidation. For this
reason certain soils and rubbish heaps--for instance, lime rubbish (in
the presence of a base)--lime contain a more or less considerable amount
of nitre. One of these nitres--sodium nitrate--is extracted from the
earth in large quantities in Chili, where it was probably formed by the
oxidation of animal refuse. This kind of nitre is employed in practice
for the manufacture of nitric acid and the other oxygen compounds of
nitrogen. Nitric acid is obtained from _Chili saltpetre_ by heating it
with _sulphuric acid_. The hydrogen of the sulphuric acid replaces the
sodium in the nitre. The sulphuric acid then forms either an acid salt,
NaHSO_{4}, or a normal salt, Na_{2}SO_{4}, whilst nitric acid is formed
from the nitre and is volatilised. The decomposition is expressed by
the equations: (1) NaNO_{3} + H_{2}SO_{4} = HNO_{3} + NaHSO_{4}, if the
acid salt be formed, and (2) 2NaNO_{3} + H_{2}SO_{4} = Na_{2}SO_{4}
+ 2HNO_{3}, if the normal sodium sulphate is formed. With an excess
of sulphuric acid, at a moderate heat, and at the commencement of the
reaction, the decomposition proceeds according to the first equation;
and on further heating with a sufficient amount of nitre according to
the second, because the acid salt NaHSO_{4} itself acts like an acid
(its hydrogen being replaceable as in acids), according to the equation
NaNO_{3} + NaHSO_{4} = Na_{2}SO_{4} + HNO_{3}.
[Illustration: FIG. 47.--Method of preparing nitric acid on a large
scale. A cast-iron retort, C, is fixed into the furnace, and heated
by the fire, B. The flame and products of combustion are at first led
along the flue, M (in order to heat the receivers), and afterwards into
L. The retort is charged with Chili saltpetre and sulphuric acid, and
the cover is luted on with clay and gypsum. A clay tube, _a_, is fixed
into the neck of the retort (in order to prevent the nitric acid from
corroding the cast iron), and a bent glass tube, D, is luted on to it.
This tube carries the vapours into a series of earthenware receivers, E.
Nitric acid mixed with sulphuric acid collects in the first. The purest
nitric acid is procured from the second, whilst that which condenses in
the third receiver contains hydrochloric acid, and that in the fourth
nitrous oxide. Water is poured into the last receiver in order to
condense the residual vapours.]
The sulphuric acid, as it is said, here displaces the nitric acid from
its compound with the base.[29 bis] Thus, in the reaction of sulphuric
acid on nitre there is formed a non-volatile salt of sulphuric acid,
which remains, together with an excess of this acid, in the distilling
apparatus, and nitric acid, which is converted into vapour, and may be
condensed, because it is a liquid and volatile substance. On a small
scale, this reaction may be carried on in a glass retort with a glass
condenser. On a large scale, in chemical works, the process is exactly
similar, only iron retorts are employed for holding the mixture of
nitre and sulphuric acid, and earthenware three-necked bottles are used
instead of a condenser,[30] as shown in fig. 47.
[29 bis] This often gives rise to the supposition that sulphuric acid
possesses a considerable degree of affinity or energy compared
with nitric acid, but we shall afterwards see that the idea of the
relative degree of affinity of acids and bases is, in many cases,
exceedingly unbiassed; it need not be accepted so long as it is
possible to explain the observed phenomena without admitting any
supposition whatever of the degree of the force of affinity,
because the latter cannot be measured. The action of sulphuric
acid upon nitre may be explained by the fact alone that the
resultant nitric acid is volatile. The nitric acid is the only one
of all the substances partaking in the reaction which is able to
pass into vapour; it alone is volatile, while the remainder are
non-volatile, or, more strictly speaking, exceedingly difficultly
volatile substances. Let us imagine that the sulphuric acid is
only able to set free a small quantity of nitric acid from its
salt, and this will suffice to explain the decomposition of the
whole of the nitre by the sulphuric acid, because once the nitric
acid is separated it passes into vapour when heated, and passes
away from the sphere of action of the remaining substances; then
the free sulphuric acid will set free a fresh small quantity of
nitric acid, and so on until it drives off the entire quantity. It
is evident that, in this explanation, it is essential that the
sulphuric acid should be in excess (although not greatly)
throughout the reaction; according to the equation expressing the
reaction, 98 parts of sulphuric acid are required per 85 parts of
Chili nitre; but if this proportion be maintained in practice the
nitric acid is not all disengaged by the sulphuric acid; an excess
of the latter must be taken, and generally 80 parts of Chili nitre
are taken per 98 parts of acid, so that a portion of the sulphuric
acid remains free to the very end of the reaction.
[30] It must be observed that sulphuric acid, at least when undiluted
(60° Baumé), corrodes cast iron with difficulty, so that the
acid may be heated in cast-iron retorts. Nevertheless, both
sulphuric and nitric acids have a certain action on cast iron,
and therefore the acid obtained will contain traces of iron. In
practice sodium nitrate (Chili saltpetre) is usually employed
because it is cheaper, but in the laboratory it is best to take
potassium nitrate, because it is purer and does not froth up so
much as sodium nitrate when heated with sulphuric acid. In the
action of an excess of sulphuric acid on nitre and nitric acid
a portion of the latter is decomposed, forming lower oxides of
nitrogen, which are dissolved in the nitric acid. A portion
of the sulphuric acid itself is also carried over as spray by
the vapours of the nitric acid. Hence sulphuric acid occurs
as an impurity in commercial nitric acid. A certain amount of
hydrochloric acid will also be found to be present in it, because
sodium chloride is generally found as an impurity in nitre, and
under the action of sulphuric acid it forms hydrochloric acid.
Commercial acid further contains a considerable excess of water
above that necessary for the formation of the hydrate, because
water is first poured into the earthenware vessels employed for
condensing the nitric acid in order to facilitate its cooling and
condensation. Further, the acid of composition HNO_{3} decomposes
with great ease, with the evolution of oxides of nitrogen. Thus
the commercial acid contains a great number of impurities, and
is frequently purified in the following manner:--Lead nitrate is
first added to the acid because it forms non-volatile and almost
insoluble (precipitated) substances with the free sulphuric
and hydrochloric acids, and liberates nitric acid in so doing,
according to the equations Pb(NO_{3})_{2} + 2HCl = PbCl_{2} +
2NHO_{3} and Pb(NO_{3})_{2} + H_{2}SO_{4} = PbSO_{4} + 2NHO_{3}.
Potassium chromate is then added to the impure nitric acid,
by which means oxygen is liberated from the chromic acid, and
this oxygen, at the moment of its evolution, oxidises the lower
oxides of nitrogen and converts them into nitric acid. A pure
nitric acid, containing no impurities other than water, may be
then obtained by carefully distilling the acid, treated as above
described, and particularly if only the middle portions of the
distillate are collected. Such acid should give no precipitate,
either with a solution of barium chloride (a precipitate shows
the presence of sulphuric acid) or with a solution of silver
nitrate (a precipitate shows the presence of hydrochloric acid),
nor should it, after being diluted with water, give a coloration
with starch containing potassium iodide (a coloration shows the
admixture of other oxides of nitrogen). The oxides of nitrogen
may be most easily removed from impure nitric acid by heating
for a certain time with a small quantity of pure charcoal. By
the action of nitric acid on the charcoal carbonic anhydride
is evolved, which carries off the lower oxides of nitrogen. On
redistilling, pure acid is obtained. The oxides of nitrogen
occurring in solution may also be removed by passing air through
the nitric acid.
Nitric acid so obtained always contains water. It is extremely difficult
to deprive it of all the admixed water without destroying a portion of
the acid itself and partially converting it into lower oxides, because
without the presence of an excess of water it is very unstable. When
rapidly distilled a portion is decomposed, and there are obtained free
oxygen and lower oxides of nitrogen, which, together with the water,
remain in solution with the nitric acid. Therefore it is necessary to
work with great care in order to obtain a pure hydrate of nitric acid,
HNO_{3}, and especially to mix the nitric acid obtained from nitre, as
above described, with sulphuric acid, which takes up the water, and
to distil it at the lowest possible temperature--that is, by placing
the retort holding the mixture in a water or oil bath and carefully
heating it. The first portion of the nitric acid thus distilled boils
at 86°, has a specific gravity at 15° of 1·526, and solidifies at -50°;
it is very unstable at higher temperatures. This is the normal hydrate,
HNO_{3}, which corresponds with the salts, NMO_{3}, of nitric acid.
When diluted with water nitric acid presents a higher boiling point,
not only as compared with that of the nitric acid itself, but also with
that of water; so that, if very dilute nitric acid be distilled, the
first portions passing over will consist of almost pure water, until
the boiling point in the vapours reaches 121°. At this temperature a
compound of nitric acid with water, containing about 70 p.c. of nitric
acid,[31] distils over; its specific gravity at 15° = 1·421. If the
solution contain less than 25 p.c. of water, then, the specific gravity
of the solution being above 1·44, HNO_{3} evaporates off and fumes in
the air, forming the above hydrate, whose vapour tension is less than
that of water. Such solutions form _fuming nitric acid_. On distilling
it gives monohydrated acid,[32] HNO_{3}; it is a hydrate boiling at
121°, so that it is obtained from both weak and strong solutions. Fuming
nitric acid, under the action not only of organic substances, but
even of heat, loses a portion of its oxygen, forming lower oxides of
nitrogen, which impart a _red-brown colour_ to it;[33] the pure acid is
colourless.
[31] Dalton, Smith, Bineau, and others considered that the hydrate of
constant boiling point (see Chapter I., Note 60) for nitric
acid was the compound 2HNO_{3},3H_{2}O, but Roscoe showed that
its composition changes with a variation of the pressure and
temperature under which the distillation proceeds. Thus, at
a pressure of 1 atmosphere the solution of constant boiling
point contains 68·6 p.c., and at one-tenth of an atmosphere
66·8 p.c. Judging from what has been said concerning solutions
of hydrochloric acid, and from the variation of specific
gravity, I think that the comparatively large decrease in the
tensions of the vapours depends on the formation of a hydrate,
NHO_{3},2H_{2}O (= 63·6 p.c.). Such a hydrate may be expressed by
N(HO)_{5}, that is, as NH_{4}(HO), in which all the equivalents
of hydrogen are replaced by hydroxyl. The constant boiling point
will then be the temperature of the decomposition of this hydrate.
The variation of the specific gravity at 15° from water (_p_
= 0) to the hydrate NHO_{3},5H_{2}O (41·2 p.c. HNO_{3}) is
expressed by _s_ = 9992 + 57·4_p_ + 0·16_p^2_, if water = 10,000
at 4°. For example, when _p_ = 30 p.c., _s_ = 11,860. For more
concentrated solutions, at least, the above-mentioned hydrate,
HNO_{3},2H_{2}O, must be taken, up to which the specific gravity
_s_ = 9570 + 84·18_p_-0·240_p_^2; but perhaps (since the results
of observations of the specific gravity of the solutions are not
in sufficient agreement to arrive at a conclusion) the hydrate
HNO_{3},3H_{2}O should be recognised, as is indicated by many
nitrates (Al, Mg, Co, &c.), which crystallise with this amount
of water of crystallisation. From HNO_{3},2H_{2}O to HNO_{3}
the specific gravity of the solutions (at 15°) _s_ = 10,652 +
62·08_p_-0·160_p_^2. The hydrate HNO_{3},2H_{2}O is recognised by
Berthelot on the basis of the thermochemical data for solutions
of nitric acid, because on approaching to this composition there
is a rapid change in the amount of heat evolved by mixing nitric
acid with water. Pickering (1892) by refrigeration obtained
the crystalline hydrates: HNO_{3},H_{2}O, melting at -37° and
HNO_{3},3H_{2}O, melting at -18°. A more detailed study of the
reactions of hydrated nitric acid would no doubt show the
existence of change in the process and rapidity of reaction in
approaching these hydrates.
[32] The normal hydrate HNO_{3}, corresponding with the ordinary salts,
may be termed the monohydrated acid, because the anhydride
N_{2}O_{5} with water forms this normal nitric acid. In this
sense the hydrate HNO_{3},2H_{2}O is the pentahydrated acid.
[33] For technical and laboratory purposes recourse is frequently had
to _red fuming nitric acid_--that is, the normal nitric acid,
HNO_{3}, containing lower oxides of nitrogen in solution. This
acid is prepared by decomposing nitre with half its weight
of strong sulphuric acid, or by distilling nitric acid with
an excess of sulphuric acid. The normal nitric acid is first
obtained, but it partially decomposes, and gives the lower
oxidation products of nitrogen, which are dissolved by the nitric
acid, to which they impart its usual pale-brown or reddish
colour. This acid fumes in the air, from which it attracts
moisture, forming a less volatile hydrate. If carbonic anhydride
be passed through the red-brown fuming nitric acid for a long
period of time, especially if, assisted by a moderate heat, it
expels all the lower oxides, and leaves a colourless acid free
from these oxides. It is necessary, in the preparation of the red
acid, that the receivers should be kept quite cool, because it
is only when cold that nitric acid is able to dissolve a large
proportion of the oxides of nitrogen. The strong red fuming acid
has a specific gravity 1·56 at 20°, and has a suffocating smell
of the oxides of nitrogen. When the red acid is mixed with water
it turns green, and then of a bluish colour, and with an excess
of water ultimately becomes colourless. This is owing to the fact
that the oxides of nitrogen in the presence of water and nitric
acid are changed, and give coloured solutions.
Markleffsky (1892) showed that the green solutions contain
(besides HNO_{3}) HNO_{2} and N_{2}O_{4}, whilst the blue
solutions only contain HNO_{2} (_see_ Note 48).
The action of red fuming nitric acid (or a mixture with sulphuric
acid) is in many cases very powerful and rapid, and it sometimes
acts differently from pure nitric acid. Thus iron becomes
covered with a coating of oxides, and insoluble in acids; it
becomes, as is said, passive. Thus chromic acid (and potassium
dichromate) gives oxide of chromium in this red acid--that is, it
is deoxidised. This is owing to the presence of the lower oxides
of nitrogen, which are capable of being oxidised--that is, of
passing into nitric acid like the higher oxides. But, generally,
the action of fuming nitric acid, both red and colourless, is
powerfully oxidising.
Nitric acid, as an _acid hydrate_, enters into reactions of double
decomposition with bases, basic hydrates (alkalis), and with salts.
In all these cases a salt of nitric acid is obtained. An alkali and
nitric acid give water and a salt; so, also, a basic oxide with nitric
acid gives a salt and water; for instance, lime, CaO + 2HNO_{3} =
Ca(NO_{3})_{2} + H_{2}O. Many of these salts are termed nitres.[34]
The composition of the ordinary salts of nitric acid may be expressed
by the general formula M(NO_{3})_{_n_}, where M indicates a metal
replacing the hydrogen in one or several (_n_) equivalents of nitric
acid. We shall find afterwards that the atoms M of metals are equivalent
to one (K, Na, Ag) atom of hydrogen, or two (Ca, Mg, Ba), or three
(Al, In), or, in general, _n_ atoms of hydrogen. _The salts of nitric
acid_ are especially characterised by being all _soluble in water_.[35]
From the property common to all these salts of entering into double
decompositions, and owing to the volatility of nitric acid, they evolve
nitric acid when heated with sulphuric acid. They all, like the acid
itself, are capable of evolving oxygen when heated, and consequently of
acting as oxidising substances; they therefore, for instance, deflagrate
with ignited carbon, the carbon burning at the expense of the oxygen of
the salt and forming gaseous products of combustion.[36]
[34] Hydrogen is not evolved in the action of nitric acid (especially
strong) on metals, even with those metals which evolve hydrogen
under the action of other acids. This is because the hydrogen
at the moment of its separation reduces the nitric acid, with
formation of the lower oxides of nitrogen, as we shall afterwards
see.
[35] Certain basic salts of nitric acid, however (for example, the
basic salt of bismuth), are insoluble in water; whilst, on the
other hand, all the normal salts are soluble, and this forms an
exceptional phenomenon among acids, because all the ordinary
acids form insoluble salts with one or another base. Thus, for
sulphuric acid the salts of barium, lead, &c., for hydrochloric
acid the salts of silver, &c., are insoluble in water. However,
the normal salts of acetic and certain other acids are all
soluble.
[36] _Ammonium nitrate_, NH_{4}NO_{3}, is easily obtained by adding a
solution of ammonia or of ammonium carbonate to nitric acid until
it becomes neutral. On evaporating this solution, crystals of
the salt are formed which contain no water of crystallisation.
It crystallises in prisms like those formed by common nitre, and
has a refreshing taste; 100 parts of water at _t_° dissolve 54 +
0·61_t_ parts by weight of the salt. It is soluble in alcohol,
melts at 160°, and is decomposed at about 180°, forming water
and nitrous oxide, NH_{4}NO_{3} = 2H_{2}O + N_{2}O. If ammonium
nitrate be mixed with sulphuric acid, and the mixture be heated
to about the boiling point of water, then nitric acid is evolved,
and ammonium hydrogen sulphate remains in solution; but if the
mixture be heated rapidly to 16O°, then nitrous oxide is evolved.
In the first case the sulphuric acid takes up ammonia, and in the
second place water. Ammonium nitrate is employed in practice for
the artificial production of cold, because in dissolving in water
it lowers the temperature very considerably. For this purpose it
is best to take equal parts by weight of the salt and water. The
salt must first be reduced to a powder and then rapidly stirred
up in the water, when the temperature will fall from +15° to -10°,
so that the water freezes.
Ammonium nitrate absorbs ammonia, with which it forms
unstable compounds resembling compounds containing water
of crystallisation. (Divers 1872, Raoult 1873.) At -10°
NH_{4}NO_{3},2NH_{3} is formed: it is a liquid of sp. gr. 1·15,
which loses all its ammonia under the influence of heat. At +28°
NH_{4}NO_{3},NH_{3} is formed: it is a solid which easily parts
with its ammonia when heated, especially in solution.
Troost (1882) investigated the tension of the dissociation of
the compounds formed, and came to the conclusion that a definite
compound corresponding to the formula 2NH_{4}NO_{3},3NH_{3} is
formed, because the tension of dissociation remains constant
in the decomposition of such a compound at 0°. Y. Kouriloff
(1893), however, considers that the constancy of the tension of
the ammonia evolved is due to the decomposition of a saturated
solution, and not of a definite compound. During decomposition
the system is composed of a liquid and a solid; the tension
only becomes constant from the moment the solid falls down. The
composition 2NH_{4}NO_{3},3NH_{3} corresponds to a saturated
solution at 0°, and the solubility of NH_{4}NO_{3} in NH_{3}
increases with a rise of temperature.
Nitric acid also enters into double decompositions with a number of
hydrocarbons not in any way possessing alkaline characters and not
reacting with other acids. Under these circumstances the nitric acid
gives water and a new substance termed a _nitro-compound_. The chemical
character of the nitro-compound is the same as that of the original
substance; for example, if an indifferent substance be taken, then the
nitro compound obtained from it will also be indifferent; if an acid
be taken, then an acid is obtained also.[36 bis] Benzene, C_{6}H_{6},
for instance, acts according to the equation C_{6}H_{6} + HNO_{3} =
H_{2}O + C_{6}H_{5}NO_{2}. Nitrobenzene is produced. The substance
taken, C_{6}H_{6}, is a liquid hydrocarbon having a faint tarry smell,
boiling at 80°, and lighter than water; by the action of nitric acid
nitrobenzene is obtained, which is a substance boiling at about 210°,
heavier than water, and having an almond-like odour: it is employed in
large quantities for the preparation of aniline and aniline dyes.[37]
As the nitro-compounds contain both combustible elements (hydrogen
and carbon), as well as oxygen in unstable combination with nitrogen,
in the form of the radicle NO_{2} of nitric acid, they decompose
with an explosion when ignited or even struck, owing to the pressure
of the vapours and gases formed--free nitrogen, carbonic anhydride,
CO_{2}, carbonic oxide, CO, and aqueous vapour. In the explosion of
nitro compounds[37 bis] much heat is evolved, as in the combustion of
gunpowder or detonating gas, and in this case the force of explosion in
a closed space is great, because from a solid or liquid nitro-compound
occupying a small space there proceed vapours and gases whose elasticity
is great not only from the small space in which they are formed, but
owing to the high temperature corresponding to the combustion of the
nitro-compound.[38]
[36 bis] This is explained by saying that in true nitro-compounds
the residue of nitric acid NO_{2} takes the place of the hydrogen
in the hydrocarbon group. For example, if C_{6}H_{5}OH be given,
then C_{6}H_{4}(NO_{2})OH will be a true nitro-compound having the
radical properties of C_{6}H_{5}OH. If, on the other hand, the
NO_{2} replace the hydrogen of the aqueous radicle
(C_{6}H_{5}ONO_{2}), then the chemical character varies, as in the
passage of KOH into KONO_{2} (nitre) (_see_ Note 37 and Organic
Chemistry).
[37] The compound ethers of nitric acid in which the hydrogen of the
aqueous radicle (OH) is replaced by the residue of nitric acid
(NO_{2}) are frequently called nitro-compounds. But in their
chemical character they differ from true nitro-compounds (for
details _see_ Organic Chemistry) and do not burn like them.
The action of nitric acid on cellulose, C_{6}H_{10}O_{5}, is an
example. This substance, which forms the outer coating of all
plant cells, occurs in an almost pure state in cotton, in common
writing-paper, and in flax, &c.; under the action of nitric
acid it forms water and nitrocellulose (like water and KNO_{3}
from KHO), which, although it has the same appearance as the
cotton originally taken, differs from it entirely in properties.
It explodes when struck, bursts into flame very easily under
the action of sparks, and acts like gunpowder, whence its name
of pyroxylin, or gun-cotton. The composition of gun-cotton is
C_{6}H_{7}N_{3}O_{11} = C_{6}H_{10}O_{5} + 3NHO_{3}-3H_{2}O.
The proportion of the group NO_{2} in nitrocellulose may be
decreased by limiting the action of the nitric acid and compounds
obtained with different properties; for instance, the (impure)
well-known _collodion cotton_, containing from 11 to 12 per
cent. of nitrogen, and _pyro-collodion_ (Mendeléeff, 1890),
containing 12·4 per cent. of nitrogen. Both these products are
soluble in a mixture of alcohol and ether (in collodion a portion
of the substance is soluble in alcohol), and the solution when
evaporated gives a transparent film, which is insoluble in water.
A solution of collodion is employed in medicine for covering
wounds, and in wet-plate photography for giving on glass an even
coating of a substance into which the various reagents employed
in the process are introduced. Extremely fine threads (obtained
by forcing a gelatinous mixture of collodion, ether, and alcohol
through capillary tubes in water) of collodion form artificial
silk.
[37 bis] The property possessed by nitroglycerin (occurring in
dynamite), nitrocellulose, and the other nitro-compounds, of
burning with an explosion, and their employment for smokeless
powder and as explosives in general, depends on the reasons in
virtue of which a mixture of nitre and charcoal deflagrates and
explodes; in both cases the elements of the nitric acid occurring
in the compound are decomposed, the oxygen in burning unites with
the carbon, and the nitrogen is set free; thus a very large volume
of gaseous substances (nitrogen and oxides of carbon) is rapidly
formed from the solid substances originally taken. These gases
occupy an incomparably larger volume than the original substance,
and therefore produce a powerful pressure and explosion. It is
evident that in exploding with the development of heat (that is,
in decomposing, not with the absorption of energy, as is generally
the case, but with the evolution of energy) the nitro-compounds
form stores of energy which are easily set free, and that
consequently their elements occur in a state of particularly
energetic motion, which is especially strong in the group NO_{2}:
this group is common to all nitro-compounds, and all the oxygen
compounds of nitrogen are unstable, easily decomposable, and (Note
29) absorb heat in their formation. On the other hand, the
nitro-compounds are instructive as an example and proof of the
fact that the elements and groups forming compounds are united in
definite order in the molecules of a compound. A blow, concussion,
or rise of temperature is necessary to bring the combustible
elements C and H into the most intimate contact with NO_{2}, and
to distribute the elements in a new order in new compounds.
As regards the composition of the nitro-compounds, it will be seen
that the hydrogen of a given substance is replaced by the complex
group NO_{2} of the nitric acid. The same is observed in the
passage of alkalis into nitrates, so that the reactions of
substitution of nitric acid--that is, the formation of salts and
nitro-compounds--may be expressed in the following manner. In
these cases the hydrogen is replaced by the so-called _radicle of
nitric acid_ NO_{2}, as is evident from the following table:--
{Caustic potash KHO.
{Nitre K(NO_{2})O.
{Hydrate of lime CaH_{2}O_{2}.
{Calcium nitrate Ca(NO_{2})_{2}O_{2}.
{Glycerin C_{3}H_{5}H_{3}O_{3}.
{Nitroglycerin C_{3}H_{5}(NO_{2})_{3}O_{3}.
{Phenol C_{6}H_{5}OH.
{Picric acid C_{6}H_{2}(NO_{2})_{3}OH, &c.
The difference between the salts formed by nitric acid and the
nitro-compounds consists in the fact that nitric acid is very
easily separated from the salts of nitric acid by means of
sulphuric acid (that is, by a method of double saline
decomposition), whilst nitric acid is not displaced by sulphuric
acid from true nitro-compounds; for instance, nitrobenzene,
C_{6}H_{5}·NO_{2}. As nitro-compounds are formed exclusively from
hydrocarbons, they are described with them in organic chemistry.
The group NO_{2} of nitro-compounds in many cases (like all the
oxidised compounds of nitrogen) passes into the ammonia group or
into the ammonia radicle NH_{2}. This requires the action of
reducing substances evolving hydrogen: RNO_{2} + 6H = RNH_{3} +
2H_{2}O. Thus Zinin converted nitrobenzene, C_{6}H_{5}·NO_{2},
into aniline, C_{6}H_{5}·NH_{2}, by the action of hydrogen
sulphide.
Admitting the existence of the group NO_{2}, as replacing hydrogen
in various compounds, then nitric acid may be considered as water
in which half the hydrogen is replaced by the radical of nitric
acid. In this sense nitric acid is nitro-water, NO_{2}OH, and its
anhydride dinitro-water, (NO_{2})_{2}O. In nitric acid the radical
of nitric acid is combined with hydroxyl, just as in nitrobenzene
it is combined with the radical of benzene.
It should here be remarked that the group NO_{3} may be recognised
in the salts of nitric acid, because the salts have the
composition M(NO_{3})_{n}, just as the metallic chlorides have the
composition MCl_{n}. But the group NO_{3} does not form any other
compounds beyond the salts, and therefore it should he considered
as hydroxyl, HO, in which H is replaced by NO_{2}.
[38] The nitro-compounds play a very important part in mining and
artillery. Detailed accounts of them must be looked for in
special works, among which the works of A. R. Shuliachenke
and T. M. Chelletsoff occupy an important place in the
Russian literature on this subject, although historically the
scientific works of Abel in England and Berthelot in France
stand pre-eminent. The latter elucidated much in connection
with explosive compounds by a series of both experimental and
theoretical researches. Among explosives a particularly important
place from a practical point of view is occupied by ordinary or
black gunpowder (Chapter XIII., Note 16), fulminating mercury
(Chapter XVI., Note 26), the different forms of gun-cotton
(Chapter VI., Note 37), and nitro-glycerine (Chapter VIII.,
Note 45, and Chapter XII., Note 33). The latter when mixed
with solid pulverulent substances, like magnesia, tripoli,
&c., forms dynamite, which is so largely used in quarries and
mines in driving tunnels, &c. We may add that the simplest true
nitro-compound, or marsh gas, CH_{4}, in which all the hydrogens
are replaced by NO_{2} groups has been obtained by L. N.
Shishkoff, C(NO_{2})_{4}, as well as nitroform, CH(NO_{2})_{3}.
[Illustration: FIG. 48.--The method of decomposition of nitrous
anhydride, also applicable to the other oxides of nitrogen, and to
their analysis. NO_{2} is generated from nitrate of lead in the retort
A. Nitric acid and other less volatile products are condensed in B.
The tube C C contains copper, and is heated from below. Undecomposed
volatile products (if any are formed) are condensed in D, which is
cooled. If the decomposition be incomplete, brown fumes make their
appearance in this receiver. The gaseous nitrogen is collected in the
cylinder E.]
If the vapour of nitric acid is passed through an even moderately
heated glass tube, the formation of dark-brown fumes of the lower
oxides of nitrogen and the separation of free oxygen may be observed,
2NHO_{3} = H_{2}O + 2NO_{2} + O. The decomposition is complete at a
white heat--that is, nitrogen is formed, 2NHO_{3} = H_{2}O + N_{2} +
O_{5}. Hence it is easily understood that nitric acid may part with its
oxygen to a number of substances capable of being oxidised.[39] It is
consequently an _oxidising agent_. Charcoal, as we have already seen,
burns in nitric acid; phosphorus, sulphur, iodine, and the majority of
metals also decompose nitric acid, some on heating and others even at
the ordinary temperature: the substances taken are oxidised and the
nitric acid is deoxidised, yielding compounds containing less oxygen.
Only a few metals, such as gold and platinum, do not act on nitric acid,
but the majority decompose it; in so doing, an oxide of the metal is
formed, which, if it has the character of a base, acts on the remaining
nitric acid; hence, with the majority of metals the result of the
reaction is usually not an oxide of the metal, but the corresponding
salt of nitric acid, and, at the same time, one of the lower oxides of
nitrogen. The resulting salts of the metals are soluble, and hence it
is said that nitric acid _dissolves_ nearly all metals.[40] This case
is termed the solution of metals by acids, although it is not a case
of simple solution, but a complex chemical change of the substances
taken. When treated with this acid, those metals whose oxides do not
combine with nitric acid yield the oxide itself, and not a salt; for
example, tin acts in this manner on nitric acid, forming a hydrated
oxide, SnH_{2}O_{3}, which is obtained in the form of a white powder,
Sn + 4NHO_{3} = H_{2}SnO_{3} + 4NO_{2} + H_{2}O. Silver is able to
take up still more oxygen, and to convert a large portion of nitric
acid into nitrous anhydride, 4Ag + 6HNO_{3} = 4AgNO_{3} + N_{2}O_{3} +
3H_{2}O. Copper takes up still more oxygen from nitric acid, converting
it into nitric oxide, and, by the action of zinc, nitric acid is able
to give up a still further quantity of nitrogen, forming nitrous oxide,
4Zn + 10NHO_{3} = 4Zn(NO_{3})_{2} + N_{2}O + 5H_{2}O.[41] Sometimes,
and especially with dilute solutions of nitric acid, the deoxidation
proceeds as far as the formation of hydroxylamine and ammonia, and
sometimes it leads to the formation of nitrogen itself. The formation of
one or other nitrogenous substance from nitric acid is determined, not
only by the nature of the reacting substances, but also by the relative
mass of water and nitric acid, and also by the temperature and pressure,
or the sum total of the conditions of reaction; and as in a given
mixture even these conditions vary (the temperature and the relative
mass vary), it not unfrequently happens that a mixture of different
products of the deoxidation of nitric acid is formed.
[39] [Illustration: FIG. 49.--Decomposition of nitrous oxide by
sodium.]
Nitric acid may be entirely decomposed by passing its vapour
over highly incandescent copper, because the oxides of nitrogen
first formed give up their oxygen to the red-hot metallic copper,
so that water and nitrogen gas alone are obtained. This forms a
means for determining the composition both of nitric acid and
of all the other compounds of nitrogen with oxygen, because
by collecting the gaseous nitrogen formed it is possible to
calculate, from its volume, its weight and consequently its
amount in a given quantity of a nitrogenous substance, and by
weighing the copper before and after the decomposition it is
possible to determine the amount of oxygen by the increase
in weight. The complete decomposition of nitric acid is also
accomplished by passing a mixture of hydrogen and nitric acid
vapours through a red-hot tube. Sodium also decomposes the oxides
of nitrogen at a red-heat, taking up all the oxygen. This method
is sometimes used for determining the composition of the oxides
of nitrogen.
[40] The application of this acid for etching copper or steel in
engraving is based on this fact. The copper is covered with a
coating of wax, resin, &c. (etching ground), on which nitric acid
does not act, and then the ground is removed in certain parts
with a needle, and the whole is washed in nitric acid. The parts
coated remain untouched, whilst the uncovered portions are eaten
into by the acid. Copper plates for etchings, aquatints, &c., are
prepared in this manner.
[41] The formation of such complex equations as the above often
presents some difficulty to the beginner. It should be observed
that if the reacting and resultant substances be known, it is
easy to form an equation for the reaction. Thus, if we wish
to form an equation expressing the reaction that nitric acid
acting on zinc gives nitrous oxide, N_{2}O, and zinc nitrate,
Zn(NO_{3})_{2}, we must reason as follows:--Nitric acid contains
hydrogen, whilst the salt and nitrous oxide do not; hence water
is formed, and therefore it is as though anhydrous nitric acid,
N_{2}O_{5}, were acting. For its conversion into nitrous oxide
it parts with four equivalents of oxygen, and hence it is able
to oxidise four equivalents of zinc and to convert it into zinc
oxide, ZnO. These four equivalents of zinc oxide require for
their conversion into the salt four more equivalents of nitric
anhydride; consequently five equivalents in all of the latter
are required, or ten equivalents of nitric acid. Thus ten
equivalents of nitric acid are necessary for four equivalents
of zinc in order to express the reaction in whole equivalents.
It must not be forgotten, however, that there are very few such
reactions which can be entirely expressed by simple equations.
The majority of equations of reactions only express the chief
and ultimate products of reaction, and thus none of the three
preceding equations express all that in reality occurs in the
action of metals on nitric acid. In no one of them is only
one oxide of nitrogen formed, but always several together or
consecutively--one after the other, according to the temperature
and strength of the acid. And this is easily intelligible. The
resulting oxide is itself capable of acting on metals and of
being deoxidised, and in the presence of the nitric acid it may
change the acid and be itself changed. The equations given must
be looked on as a systematic expression of the main features of
reactions, or as a limit towards which they tend, but to which
they only attain in the absence of disturbing influences.
Thus the action of nitric acid on metals consists in their being
oxidised, whilst the acid itself is converted, according to the
temperature, concentration in which it is taken, and the nature of the
metal, &c., into lower oxides, ammonia, or even into nitrogen.[42] Many
compounds are oxidised by nitric acid like metals and other elements;
for instance, lower oxides are converted into higher oxides. Thus,
arsenious acid is converted into arsenic acid, suboxide of iron into
oxide, sulphurous acid into sulphuric acid, the sulphides of the metals,
M_{2}S, into sulphates, M_{2}SO_{4}, &c.; in a word, nitric acid brings
about oxidation, its oxygen is taken up and transferred to many other
substances. Certain substances are oxidised by strong nitric acid so
rapidly and with so great an evolution of heat that they deflagrate
and burst into flame. Thus turpentine, C_{10}H_{16}, bursts into
flame when poured into fuming nitric acid. In virtue of its oxidising
property, nitric acid _removes the hydrogen_ from many substances. Thus
it decomposes hydriodic acid, separating the iodine and forming water;
and if fuming nitric acid be poured into a flask containing gaseous
hydriodic acid, then a rapid reaction takes place, accompanied by flame
and the separation of violet vapours of iodine and brown fumes of oxides
of nitrogen.[43]
[42] Montemartini endeavours to show that the products evolved in the
action of nitric acid upon metals (and their amount) is in direct
connection with both the concentration of the acid and the
capacity of the metals to decompose water. Those metals which
only decompose water at a high temperature give, under the action
of nitric acid, NO_{2}, N_{2}O_{4}, and NO; whilst those metals
which decompose water at a lower temperature give, besides the
above products, N_{2}O, N, and NH_{3}; and, lastly, the metals
which decompose water at the ordinary temperature also evolve
hydrogen. It is observed that concentrated nitric acid oxidises
many metals with much greater difficulty than when diluted
with water; iron, copper, and tin are very easily oxidised by
dilute nitric acid, but remain unaltered under the influence of
monohydrated nitric acid or of the pure hydrate NHO_{3}. Nitric
acid diluted with a large quantity of water does not oxidise
copper, but it oxidises tin; dilute nitric acid also does not
oxidise either silver or mercury; but, on the addition of nitrous
acid, even dilute acid acts on the above metals. This naturally
depends on the smaller stability of nitrous acid, and on the fact
that after the commencement of the action the nitric acid is
itself converted into nitrous acid, which continues to act on the
silver and mercury. Veley (Oxford 1891) made detailed researches
on the action of nitric acid upon Cu, Hg, and Bi, and showed that
nitric acid of 30 p.c. strength does not act upon these metals at
the ordinary temperature if nitrous acid (traces are destroyed by
urea) and oxidising agents such as H_{2}O_{2}, KClO_{3}, &c. be
entirely absent; but in the presence of even a small amount of
nitrous acid the metals form nitrites, which, with HNO_{3}, form
nitrates and the oxides of nitrogen, which re-form the nitrous
acid necessary for starting the reaction, because the reaction
2NO + HNO_{3} + H_{2}O = 3HNO_{2} is reversible. The above metals
are quickly dissolved in a 1 p.c. solution of nitrous acid.
Moreover, Veley observed that nitric acid is partially converted
into nitrous acid by gaseous hydrogen in the presence of the
nitrates of Cu and Pb.
[43] When nitric acid acts on many organic substances it often happens
that not only is hydrogen removed, but also oxygen is combined;
thus, for example, nitric acid converts toluene, C_{7}H_{8},
into benzoic acid, C_{7}H_{6}O_{2}. In certain cases, also, a
portion of the carbon contained in an organic substance burns at
the expense of the oxygen of the nitric acid. So, for instance,
phthalic acid, C_{8}H_{6}O_{4}, is obtained from naphthalene,
_{10}H_{8}. Thus the action of nitric acid on the hydrocarbons
is often most complex; not only does nitrification take place,
but also separation of carbon, displacement of hydrogen, and
combination of oxygen. There are few organic substances which can
withstand the action of nitric acid, and it causes fundamental
changes in a number of them. It leaves a yellow stain on the
skin, and in a large quantity causes a wound and entirely eats
away the membranes of the body. The membranes of plants are eaten
into with the greatest ease by strong nitric acid in just the
same manner. One of the most durable blue vegetable dyes employed
in dyeing tissues is _indigo_; yet it is easily _converted into a
yellow substance_ by the action of nitric acid, and small traces
of free nitric acid may be recognised by this means.
As nitric acid is very easily decomposed with the separation of oxygen,
it was for a long time supposed that it was not capable of forming the
corresponding _nitric anhydride_, N_{2}O_{5}; but Deville first and
subsequently Weber and others, discovered the methods of its formation.
Deville obtained nitric anhydride by decomposing silver nitrate by
chlorine under the influence of a moderate heat. Chlorine acts on
the above salt at a temperature of 95° (2AgNO_{3} + Cl_{2} = 2AgCl +
N_{2}O_{5} + O), and when once the reaction is started, it continues
by itself without further heating. Brown fumes are given off, which
are condensed in a tube surrounded by a freezing-mixture. A portion
condenses in this tube and a portion remains in a gaseous state. The
latter contains free oxygen. A crystalline mass and a liquid substance
are obtained in the tube; the liquid is poured off, and a current of dry
carbonic acid gas is passed through the apparatus in order to remove
all traces of volatile substances (liquid oxides of nitrogen) adhering
to the crystals of nitric anhydride. These form a voluminous mass of
rhombic crystals (density 1·64), which sometimes are of rather large
size; they melt at about 30° and distil at about 47°. In distilling,
a portion of the substance is decomposed. With water these crystals
give nitric acid. Nitric anhydride is also obtained by the action of
phosphoric anhydride, P_{2}O_{5}, on cold pure nitric acid (below
0°). During the very careful distillation of equal parts by weight
of these two substances a portion of the acid decomposes, giving a
liquid compound, H_{2}O,2N_{2}O_{5} = N_{2}O_{5},2HNO_{3}, whilst the
greater part of the nitric acid gives the anhydride according to the
equation 2NHO_{3} + P_{2}O_{5} = 2PHO_{3} + N_{2}O_{5}. On heating,
nitric anhydride decomposes with an explosion, or gradually, into nitric
peroxide and oxygen, N_{2}O_{5} = N_{2}O_{4} + O.
_Nitrogen peroxide_, N_{2}O_{4}, and _nitrogen dioxide_, NO_{2}, express
one and the same composition, but they should be distinguished like
ordinary oxygen and ozone, although in this case their mutual conversion
is more easily effected and takes place on vaporisation; also, O_{3}
loses heat in passing into O_{2}, whilst N_{2}O_{4} absorbs heat in
forming NO_{2}.
Nitric acid in acting on tin and on many organic substances (for
example, starch) gives brown vapours, consisting of a mixture of
N_{2}O_{3} and NO_{2}. A purer product is obtained by the decomposition
of lead nitrate by heat, Pb(NO_{3})_{2} = 2NO_{2} + O + PbO, when
non-volatile lead oxide, oxygen gas, and nitrogen peroxide are formed.
The latter condenses, in a well-cooled vessel, to a brown liquid, which
boils at about 22°. The purest peroxide of nitrogen, solidifying at -9°,
is obtained by mixing dry oxygen in a freezing-mixture with twice its
volume of dry nitric oxide, NO, when transparent prisms of nitrogen
peroxide are formed in the receiver: they melt into a colourless liquid
at about -10°. When the temperature of the receiver is above -9°, the
crystals melt,[44] and at 0° give a reddish yellow liquid, like that
obtained in the decomposition of lead nitrate. The vapours of nitrogen
peroxide have a characteristic odour, and at the ordinary temperature
are of a dark-brown colour, but at lower temperatures the colour of
the vapour is much fainter. When heated, especially above 50°, the
colour becomes a very dark brown, so that the vapours almost lose their
transparency.
[44] According to certain investigations, if a brown liquid is formed
from the melted crystals by beating above -9°, then they no longer
solidify at -10°, probably because a certain amount of N_{2}O_{3}
(and oxygen) is formed, and this substance remains liquid at
-30°, or it may be that the passage from 2NO_{2} into N_{2}O_{4}
is not so easily accomplished as the passage from N_{2}O_{4} into
2NO_{2}.
Liquid nitrogen peroxide (that is, a mixture of NO_{2} and
N_{2}O_{4}) is employed in admixture with hydrocarbons as an
explosive.
The causes of these peculiarities of nitrogen peroxide were not
clearly understood until Deville and Troost determined the density and
dissociation of the vapour of this substance at different temperatures,
and showed that the density varies. If the density be referred to that
of hydrogen at the same temperature and pressure, then it is found to
vary from 38 at the boiling point, or about 27°, to 23 at 135°, after
which the density remains constant up to those high temperatures at
which the oxides of nitrogen are decomposed. As on the basis of the
laws enunciated in the following chapter, the density 23 corresponds
with the compound NO_{2} (because the weight corresponding with this
molecular formula = 46, and the density referred to hydrogen as unity
is equal to half the molecular weight); therefore at temperatures above
135° the existence of nitrogen dioxide only must be recognised. It is
this gas which is of a brown colour. At a lower temperature it forms
nitrogen peroxide, N_{2}O_{4}, whose molecular weight, and therefore
density, is twice that of the dioxide. This substance, which is isomeric
with nitrogen dioxide, as ozone is isomeric with oxygen, and has twice
as great a vapour density (46 referred to hydrogen), is formed in
greater quantity the lower the temperature, and crystallises at -10°. The
reasons both of the variation of the colour of the gas (N_{2}O_{4} gives
colourless and transparent vapours, whilst those of NO_{2} are brown
and opaque) and the variation of the vapour density with the variation
of temperature are thus made quite clear; and as at the boiling point
a density 38 was obtained, therefore at that temperature the vapours
consist of a mixture of 79 parts by weight of N_{2}O_{4} with 21 parts
by weight of NO_{2}.[45] It is evident that a decomposition here takes
place the peculiarity of which consists in the fact that the product of
decomposition, NO_{2}, is polymerised (_i.e._ becomes denser, combines
with itself) at a lower temperature; that is, the reaction
N_{2}O_{4} = NO_{2} + NO_{2}
is a reversible reaction, and consequently the whole phenomenon
represents a _dissociation_ in a homogeneous gaseous medium, where the
original substance, N_{2}O_{4}, and the resultant, NO_{2}, are both
gases. The _measure of dissociation_ will be expressed if we find the
proportion of the quantity of the substance decomposed to the whole
amount of the substance. At the boiling point, therefore, the measure
of the decomposition of nitrogen peroxide will be 21 p.c.; at 135° it =
1, and at 10° it = 0; that is, the N_{2}O_{4} is not then decomposable.
Consequently the limits of dissociation here are -10° and 135° at the
atmospheric pressure.[46] Within the limits of these temperatures the
vapours of nitrogen peroxide have not a constant density, but, on the
other hand, above and below these limits definite substances exist. Thus
above 135° N_{2}O_{4} has ceased to exist and NO_{2} alone remains.
It is evident that at the ordinary temperature there is a partially
dissociated system or mixture of nitrogen peroxide, N_{2}O_{4}, and
nitrogen dioxide, NO_{2}. In the brown liquid boiling at 22° probably
a portion of the N_{2}O_{4} has already passed into NO_{2}, and it is
only the colourless liquid and crystalline substance at -10° that can be
considered as pure nitrogen peroxide.[47]
[45] Because if _x_ equal the amount by weight of N_{2}O_{4}, its
volume will = _x_/46, and the amount of NO_{2} will = 100-_x_,
and consequently its volume will = (100-_x_)/23. But the mixture,
having a density 38, will weigh 100; consequently its volume will
= 100/38. Hence _x_/46 + (100-_x_)/23 = 100/38, or _x_ = 79·O.
[46] The phenomena and laws of dissociation, which we shall consider
only in particular instances, are discussed in detail in works
on theoretical chemistry. Nevertheless, in respect to nitrogen
peroxide, as an historically important example of dissociation
in a homogeneous gaseous medium, we will cite the results of the
careful investigations (1885-1880) of E. and L. Natanson, who
determined the densities under various conditions of temperature
and pressure. The degree of dissociation, expressed as above (it
may also he expressed otherwise--for example, by the ratio of
the quantity of substance decomposed to that unaltered), proves
to increase at all temperatures as the pressure diminishes,
which would he expected for a homogeneous gaseous medium, as a
decreasing pressure aids the formation of the lightest product of
dissociation (that having the least density or largest volume).
Thus, in the Natansons' experiments the degree of dissociation
at 0° increases from 10 p.c. to 30 p.c., with a decrease of
pressure of from 251 to 38 mm.; at 49°·7 it increases from 49
p.c. to 93 p.c., with a fall of pressure of from 498 to 27 mm.,
and at 100° it increases from 89·2 p.c. to 99·7 p.c., with a
fall of pressure of from 732·5 to 11·7 mm. At 130° and 150° the
decomposition is complete--that is, only NO_{2} remains at the
low pressures (less than the atmospheric) at which the Natansons
made their determinations; but it is probable that at higher
pressures (of several atmospheres) molecules of N_{2}O_{4} would
still be formed, and it would be exceedingly interesting to trace
the phenomena under the conditions of both very considerable
pressures and of relatively large volumes.
[47] Liquid nitrogen peroxide is said by Geuther to boil at 22°-26°,
and to have a sp. gr. at 0° = 1·494 and at 15° = 1·474. It
is evident that, in the liquid as in the gaseous state, the
variation of density with the temperature depends, not only
on physical, but also on chemical changes, as the amount of
N_{2}O_{4} decreases and the amount of NO_{2} increases with
the temperature, and they (as polymeric substances) should
have different densities, as we find, for instance, in the
hydrocarbons C_{5}H_{10} and C_{10}H_{20}.
It may not be superfluous to mention here that the measurement
of the specific heat of a mixture of the vapours of N_{2}O_{4}
and NO_{2} enabled Berthelot to determine that the transformation
of 2NO_{2} into N_{2}O_{4} is accompanied by the evolution of
about 13,000 units of heat, and as the reaction proceeds with
equal facility in either direction, it will be exothermal in the
one direction and endothermal in the other; and this clearly
demonstrates the possibility of reactions taking place in either
direction, although, as a rule, reactions evolving heat proceed
with greater ease.
The above explains the action of nitrogen peroxide on water at low
temperatures. N_{2}O_{4} then acts on water like a mixture of the
anhydrides of nitrous and nitric acids. The first, N_{2}O_{3}, may
be looked on as water in which each of the two atoms of hydrogen is
replaced by the radicle NO, while in the second each hydrogen is
replaced by the radicle NO_{2}, proper to nitric acid; and in nitrogen
peroxide one atom of the hydrogen of water is replaced by NO and the
other by NO_{2}, as is seen from the formulæ--
H} NO} NO } NO_{2}}
H} O; NO} O; NO_{2}} O; NO_{2}} O;
or H_{2}O; N_{2}O{3}; N_{2}O_{4}; N_{2}O_{5}.
In fact, nitrogen peroxide at low temperatures gives with water (ice)
both nitric, HNO_{3}, and nitrous, HNO_{2}, acids. The latter, as
we shall afterwards see, splits up into water and the anhydride,
N_{2}O_{3}. If, however, warm water act on nitrogen peroxide, only
nitric acid and monoxide of nitrogen are formed: 3NO_{2} + H_{2}O = NO +
2NHO_{3}.
Although NO_{2} is not decomposed into N and O even at 500°, still
in many cases it acts as an oxidising agent. Thus, for instance, it
oxidises mercury, converting it into mercurous nitrate, 2NO_{2} + Hg
= HgNO_{3} + NO, being itself deoxidised into nitric oxide, into which
the dioxide in many other instances passes, and from which it is easily
formed.[48]
[48] Nitric acid of sp. gr. 1·51 in dissolving nitrogen peroxide becomes
brown, whilst nitric acid of sp. gr. 1·32 is coloured greenish
blue, and acid of sp. gr. below 1·15 remains colourless after
absorbing nitrogen peroxide (Note 33).
_Nitrous anhydride_, N_{2}O_{3}, corresponds[49] to nitrous acid,
NHO_{2}, which forms a series of salts, the nitrites--for example,
the sodium salt NaNO_{2}, the potassium salt KNO_{2}, the ammonium
salt (NH_{4})NO_{2},[50] the silver salt AgNO_{2},[51] &c. Neither
the anhydride nor the hydrate of the acid is known in a perfectly
pure state. The anhydride has only been obtained as a very unstable
substance, and has not yet been fully investigated; and on attempting
to obtain the acid NHO_{2} from its salts, it always gives water and
the anhydride, whilst the latter, as an intermediate oxide, partially
or wholly splits up into NO + NO_{2}. But the salts of nitrous acid are
distinguished for their great stability. Potassium nitrate, KNO_{3}, may
be converted into potassium nitrite by depriving it of a portion of
its oxygen; for instance, by fusing it (at not too high a temperature)
with metals, such as lead, KNO_{3} + Pb = KNO_{2} + PbO.[51 bis]
The resultant salt is soluble in water, whilst the oxide of lead is
insoluble. With sulphuric and other acids the solution of potassium
nitrite[52] immediately evolves a brown gas, nitrous anhydride:
2KNO_{2} + H_{2}SO_{4} = K_{2}SO_{4} + N_{2}O_{3} + H_{2}O. The same
gas (N_{2}O_{3}) is obtained by passing nitric oxide at 0° through
liquid peroxide of nitrogen,[53] or by heating starch with nitric acid
of sp. gr. 1·3. At a very low temperature it condenses into a blue
liquid boiling below 0°,[54] but then partially decomposing into NO +
NO_{2}. Nitrous anhydride possesses a remarkable capacity for oxidising.
Ignited bodies burn in it, nitric acid absorbs it, and then acquires
the property of acting on silver and other metals, even when diluted.
_Potassium iodide_ is oxidised by this gas just as it is by ozone (and
by peroxide of hydrogen, chromic and other acids, but not by dilute
nitric acid nor by sulphuric acid), with the _separation of iodine_.
This iodine may he recognised (_see_ Ozone, Chapter IV.) by its turning
starch blue. Very small traces of nitrites may be easily detected by
this method. If, for example, starch and potassium iodide are added
to a solution of potassium nitrite (at first there will be no change,
there being no free nitrous acid), and then sulphuric acid be added, the
nitrous acid (or its anhydride) immediately set free liberates iodine,
which produces a blue colour with the starch. Nitric acid does not act
in this manner, but in the presence of zinc the coloration takes place,
which proves the formation of nitrous acid in the deoxidation of nitric
acid.[55] Nitrous acid acts directly on ammonia, forming nitrogen and
water, HNO_{2} + NH_{3} = N_{2} + 2H_{2}O.[56]
[49] Nitrogen peroxide as a mixed substance has no corresponding
independent salts, but Sabatier and Senderens (1892) showed that
under certain conditions NO_{2} combines directly with some
metals--for instance, copper and cobalt--forming Cu_{2}NO_{2} and
CoNO_{2} as dark brown powders, which do not, however, exhibit
the reactions of salts. Thus by passing gaseous nitrogen dioxide
over freshly reduced (from the oxides by heating with hydrogen)
copper at 25°-30°, Cu_{2}NO_{2} is directly formed. With water
it partly gives off NO_{2} and partly forms nitrite of copper,
leaving metallic copper and its suboxide. The nature of these
compounds has not yet been sufficiently investigated.
[50] Ammonium nitrite may be easily obtained in solution by a similar
method of double decomposition (for instance, of the barium salt
with ammonium sulphate) to the other salts of nitrous acid, but
it decomposes with great ease when evaporated, with evolution
of gaseous nitrogen, as already mentioned (Chapter V.) If the
solution, however, be evaporated at the ordinary temperature
under the receiver of an air-pump, a solid saline mass is
obtained, which is easily decomposed when heated. The dry salt
even decomposes with an explosion when struck, or when heated to
about 70°--NH_{4}NO_{2} = 2H_{2}O + N_{2}. It is also formed by
the action of aqueous ammonia on a mixture of nitric oxide and
oxygen, or by the action of ozone on ammonia, and in many other
instances. Zörensen (1894) prepared NH_{4}NO_{2} by the action
of a mixture of N_{2}O_{3} and other oxides of nitrogen on lumps
of ammonium carbonate, extracting the nitrite of ammonium formed
with absolute alcohol, and precipitating it from this solution
by ether. This salt is crystalline, dissolves in water with
absorption of heat, and attracts moisture from the air. The solid
salt and its concentrated solutions decompose with an explosion
when heated to 50°-80°, especially in the presence of traces
of foreign acids. Decomposition also proceeds at the ordinary
temperature, but more slowly; and in order to preserve the salt
it should be covered with a layer of pure dry ether.
[51] Silver nitrite, AgNO_{2}, is obtained as a very slightly soluble
substance, as a precipitate, on mixing solutions of silver
nitrate, AgNO_{3}, and potassium nitrite, KNO_{2}. It is soluble
in a large volume of water, and this is taken advantage of
to free it from silver oxide, which is also present in the
precipitate, owing to the fact that potassium nitrite always
contains a certain amount of oxide, which with water gives the
hydroxide, forming oxide of silver with silver nitrate. The
solution of silver nitrite gives, by double decomposition with
metallic chlorides (for instance, barium chloride), insoluble
silver chloride and the nitrite of the metal taken (in this case,
barium nitrite, Ba(NO_{2})_{2}).
[51 bis] Leroy (1889) obtained KNO_{2} by mixing powdered KNO_{3} with
BaS, igniting the mixture in a crucible and washing the fused
salts; BaSO_{4} is then left as an insoluble residue, and KNO_{2}
passes into solution: 4KNO_{3} + BaS = 4KNO_{2} + BaSO_{4}.
[52] Probably potassium nitrite, KNO_{2}, when strongly heated,
especially with metallic oxides, evolves N and O, and gives
potassium oxide, K_{2}O, because nitre is liable to such a
decomposition; but it has, as yet, been but little investigated.
[53] There are many researches which lead to the conclusion that the
reaction N_{2}O_{3} = NO_{2}-NO is reversible, _i.e._ resembles
the conversion of N_{2}O_{4} into NO_{2}. The brown colour of the
fumes of N_{2}O_{3} is due to the formation of NO_{2}.
If nitrogen peroxide be cooled to -20°, and half its weight
of water be added to it drop by drop, then the peroxide is
decomposed, as we have already said, into nitrous and nitric
acids; the former does not then remain as a hydrate, but
straightway passes into the anhydride, and, hence, if the
resultant liquid be slightly warmed vapours of nitrous anhydride,
N_{2}O_{3}, are evolved, and condense into a blue liquid, as
Fritzsche showed. This method of preparing nitrous anhydride
apparently gives the purest product, but it easily dissociates,
forming NO and NO_{2} (and therefore also nitric acid in the
presence of water).
[54] According to Thorpe, N_{2}O_{3} boils at +18°. According to
Geuther, at +3°·5, and its sp. gr. at 0° = 1·449.
[55] In its oxidising action nitrous anhydride gives nitric oxide,
N_{2}O_{3} = 2NO + O. Thus its analogy to ozone becomes still
more marked, because in ozone it is only one-third of the oxygen
that acts in oxidising; from O_{3} there is obtained O, which
acts as an oxidiser, and common oxygen O_{2}. In a physical
aspect the relation between N_{2}O_{3} and O_{3} is revealed in
the fact that both substances are of a blue colour when in the
liquid state.
[56] This reaction is taken advantage of for converting the amides,
NH_{2}R (where R is an element or a complex group) into
hydroxides, RHO. In this case NH_{2}R + NHO_{2} forms 2N
+ H_{2}O + RHO; NH_{2}, is replaced by HO, the radicle of
ammonia by the radicle of water. This reaction is employed
for transforming many nitrogenous organic substances having
the properties of amides into their corresponding hydroxides.
Thus aniline, C_{6}H_{5}·NH_{2}, which is obtained from
nitrobenzene, C_{6}H_{5}·NO_{2} (Note 37), is converted by
nitrous anhydride into phenol, C_{6}H_{5}·OH, which occurs in the
creosote extracted from coal tar. Thus the H of the benzene is
successively replaced by NO_{2}, NH_{2}, and HO; a method which
is suitable for other cases also.
As nitrous anhydride easily splits up into NO_{2} + NO, so, like NO_{2},
with warm water it gives nitric acid and nitric oxide, according to the
equation 3N_{2}O_{3} + H_{2}O = 4NO + 2NHO_{3}.
Being in a lower degree of oxidation than nitric acid, nitrous
acid and its anhydride are oxidised in solutions by many oxidising
substances--for example, by potassium permanganate--into nitric acid.[57]
[57] The action of a solution of potassium permanganate, KMnO_{4},
on nitrous acid in the presence of sulphuric acid is determined
by the fact that the higher oxide of manganese, Mn_{2}O_{7},
contained in the permanganate is converted into the lower oxide,
MnO, which as a base forms manganese sulphate, MnSO_{4}, and
the oxygen serves for the oxidation of the N_{2}O_{3} into
N_{2}O_{5}, or its hydrate. As the solution of the permanganate
is of a red colour, whilst that of manganese sulphate is almost
colourless, this reaction is clearly seen, and may be employed
for the detection and determination of nitrous acid and its salts.
_Nitric oxide_, NO.--This permanent gas[58] (that is, unliquefiable
by pressure without the aid of cold) may be obtained from all the
above-described compounds of nitrogen with oxygen. The deoxidation of
nitric acid by metals is the usual method employed for its preparation.
Dilute nitric acid (sp. gr. 1·18, but not stronger, as then N_{2}O_{3}
and NO_{2} are produced) is poured into a flask containing metallic
copper.[59] The reaction commences at the ordinary temperature. Mercury
and silver also give nitric oxide with nitric acid. In these reactions
with metals one portion of the nitric acid is employed in the oxidation
of the metal, whilst the other, and by far the greater, portion combines
with the metallic oxide so obtained, with formation of the nitrate
corresponding with the metal taken. The first action of the copper on
the nitric acid is thus expressed by the equation
2NHO_{3} + 3Cu = H_{2}O + 3CuO + 2NO.
The second reaction consists in the formation of copper nitrate--
6NHO_{3} + 3CuO = 3H_{2}O + 3Cu(NO_{3})_{2}.
[58] The absolute boiling point = -93° (_see_ Chapter II., Note 29).
[59] Kammerer proposed preparing nitric oxide, NO, by pouring a solution
of sodium nitrate over copper shavings, and adding sulphuric
acid drop by drop. The oxidation of ferrous salts by nitric acid
also gives NO. One part of strong hydrochloric acid is taken and
iron is dissolved in it (FeCl_{2}), and then an equal quantity
of hydrochloric acid and nitre is added to the solution. On
heating, nitric oxide is evolved. In the presence of an excess
of sulphuric acid and mercury the conversion of nitric acid into
nitric oxide is complete (that is, the reaction proceeds to the
end and the nitric oxide is obtained without other products), and
upon this is founded one of the methods for determining nitric
acid (in nitrometers of various kinds, described in text-books
of analytical chemistry), as the amount of NO can be easily
and accurately measured volumetrically. The amount of nitrogen
in gun-cotton, for instance, is determined by dissolving it in
sulphuric acid. Nitrous acid acts in the same manner. Upon this
property Emich (1892) founds his method for preparing pure NO.
He pours mercury into a flask, and then covers it with sulphuric
acid, in which a certain amount of NaNO_{2} or other substance
corresponding to HNO_{2} or HNO_{3} has been dissolved. The
evolution of NO proceeds at the ordinary temperature, being more
rapid as the surface of the mercury is increased (if shaken, the
reaction proceeds very rapidly). If the gas be passed over KHO,
it is obtained quite pure, because KHO does not act upon NO at
the ordinary temperature (if heated, KNO_{2} and N_{2}O or N_{2},
are formed).
Nitric oxide is a colourless gas which is only slightly soluble in
water (1/20 of a volume at the ordinary temperature). Reactions of
double decomposition in which nitric oxide readily takes part are not
known--that is to say, it is an indifferent, not a saline, oxide.
Like the other oxides of nitrogen, it is decomposed into its elements
at a red heat (starting from 900°, at 1,200° 60 per cent. give N_{2}
and 2N_{2}O_{3}, but complete decomposition into N_{2} and O_{2} only
takes place at the melting point of platinum, Emich 1892). The most
characteristic property of nitric oxide is its capacity for directly
and easily combining with oxygen (owing to the evolution of heat in
the combination). With oxygen it forms nitrous anhydride and nitrogen
peroxide, 2NO + O = N_{2}O_{3}, 2NO + O_{2} = 2NO_{2}. If nitric oxide
is mixed with oxygen and immediately shaken up with caustic potash,
it is almost entirely converted into potassium nitrite; whilst after
a certain time, when the formation of nitric peroxide has already
commenced, a mixture of potassium nitrite and nitrate is obtained.
If oxygen is passed into a bell jar filled with nitric oxide, brown
fumes of nitrous anhydride and nitric peroxide are formed, even in the
absence of moisture; these in the presence of water give, as we already
know, nitric acid and nitric oxide, so that in the presence of an
excess of water and oxygen the whole of the nitric oxide is easily and
directly converted into nitric acid. This reaction of the re-formation
of nitric acid from nitric oxide, air, and water, 2NO + H_{2}O + O_{3}
= 2HNO_{3}, is frequently made use of in practice. The experiment
showing the conversion of nitric oxide into nitric acid is very striking
and instructive. As the intermixture of the oxygen with the oxide of
nitrogen proceeds, the nitric acid formed dissolves in water, and if an
excess of oxygen has not been added the whole of the gas (nitric oxide),
being converted into HNO_{3}, is absorbed, and the water entirely fills
the bell jar previously containing the gas.[60] It is evident that
nitric oxide[61] in combining with oxygen has a strong tendency to give
the higher types of nitrogen compounds, which we see in nitric acid,
HNO_{3} or NO_{2}(OH), in nitric anhydride, N_{2}O_{5} or (NO_{2})_{2}O,
and in ammonium chloride, NH_{4}Cl. If X stand for an atom of hydrogen,
or its equivalents, chlorine, hydroxyl, &c., and if O, which is,
according to the law of substitution, equivalent to H_{2}, be indicated
by X_{2}, then the three compounds of nitrogen above named should be
considered as compounds of the type or form NX_{5}. For example, in
nitric acid X_{5} = O_{2} + (OH), where O_{2} = X_{4}, and OH = X;
whilst nitric oxide is a compound of the form NX_{2}. Hence this lower
form, like lower forms in general, strives by combination to attain to
the higher forms proper to the compounds of a given element. NX_{2}
passes consecutively into NX_{3}--namely, into N_{2}O_{3} and NHO_{2},
NX_{4} (for instance NO_{2}) and NX_{5}.
[60] This transformation of the permanent gases nitric oxide and oxygen
into liquid nitric acid in the presence of water, and with
the evolution of heat, presents a most striking instance
of liquefaction produced by the action of chemical forces.
They perform with ease the work which physical (cooling) and
mechanical (pressure) forces effect with difficulty. In this the
motion, which is so distinctively the property of the gaseous
molecules, is apparently destroyed. In other cases of chemical
action it is apparently created, arising, no doubt, from latent
energy--that is, from the internal motion of the atoms in the
molecules.
[61] Nitric oxide is capable of entering into many characteristic
combinations; it is absorbed by the solutions of many acids, for
instance, tartaric, acetic, phosphoric, sulphuric, and metallic
chlorides (for example, SbCl_{5}, BiCl_{3}, &c., with which it
forms definite compounds; Besson 1889), and also by the solutions
of many salts, especially those formed by suboxide of iron (for
instance, ferrous sulphate). In this case a brown compound is
formed which is exceedingly unstable, like all the analogous
compounds of nitric oxide. The amount of nitric oxide combined
in this manner is in atomic proportion with the amount of the
substance taken; thus ferrous sulphate, FeSO_{4}, absorbs it in
the proportion of NO to 2FeSO_{4}. Ammonia is obtained by the
action of a caustic alkali on the resultant compound, because
the oxygen of the nitric oxide and water are transferred to the
ferrous oxide, forming ferric oxide, whilst the nitrogen combines
with the hydrogen of the water. According to the investigations
of Gay (1885), the compound is formed with the evolution of
a large quantity of heat, and is easily dissociated, like a
solution of ammonia in water. It is evident that oxidising
substances (for example, potassium permanganate, KMnO_{4}, Note
57) are able to convert it into nitric acid. If the presence
of a radicle NO_{2}, composed like nitrogen peroxide, must be
recognised in the compounds of nitric acid, then a radicle NO,
having the composition of nitric oxide, may be admitted in the
compounds of nitrous acid. The compounds in which the radicle NO
is recognised are called _nitroso-compounds_. These substances
are described in Prof. Bunge's work (Kief, 1868).
As the decomposition of nitric oxide begins at temperatures above 900°,
many substances burn in it; thus, ignited phosphorus continues to burn
in nitric oxide, but sulphur and charcoal are extinguished in it. This
is due to the fact that the heat evolved in the combustion of these two
substances is insufficient for the decomposition of the nitric oxide,
whilst the heat developed by burning phosphorus suffices to produce this
decomposition. That nitric oxide really supports combustion, owing to
its being decomposed by the action of heat, is proved by the fact that
strongly ignited charcoal continues to burn in the same nitric oxide[62]
in which a feebly incandescent piece of charcoal is extinguished.
[62] A mixture of nitric oxide and hydrogen is inflammable. If a
mixture of the two gases be passed over spongy platinum the
nitrogen and hydrogen even combine, forming ammonia. A mixture
of nitric oxide with many combustible vapours and gases is very
inflammable. A very characteristic flame is obtained in burning a
mixture of nitric oxide and the vapour of the combustible carbon
bisulphide, CS_{2}. The latter substance is very volatile, so
that it is sufficient to pass the nitric oxide through a layer
of the carbon bisulphide (for instance, in a Woulfe's bottle)
in order that the gas escaping should contain a considerable
amount of the vapours of this substance. This mixture continues
to burn when ignited, and the flame emits a large quantity of
the so-called ultra-violet rays, which are capable of inducing
chemical combinations and decompositions, and therefore the flame
may be employed in photography in the absence of sufficient
daylight (magnesium light and electric light have the same
property). There are many gases (for instance, ammonia) which
when mixed with nitric oxide explode in a eudiometer.
The compounds of nitrogen with oxygen which we have so far considered
may all be prepared from nitric oxide, and may themselves be converted
into it. Thus nitric oxide stands in intimate connection with them.[63]
The passage of nitric oxide into the higher degrees of oxidation and the
converse reaction is employed in practice as a means for _transferring_
the oxygen of the air to substances capable of being oxidised. Starting
with nitric oxide, it may easily be converted, with the aid of the
oxygen of the atmosphere and water, into nitric acid, nitrous anhydride,
and nitric peroxide, and by their means employed to oxidise other
substances. In this oxidising action nitric oxide is again formed, and
it may again be converted into nitric acid, and so on continuously,
if only oxygen and water be present. Hence the fact, which at first
appears to be a paradox, that by means of a small quantity of nitric
oxide in the presence of oxygen and water it is possible to oxidise
an indefinitely large quantity of substances which cannot be directly
oxidised either by the action of the atmospheric oxygen or by the
action of nitric oxide itself. The sulphurous anhydride, SO_{2}, which
is obtained in the combustion of sulphur and in roasting many metallic
sulphides in the air is an example of this kind. In practice this gas is
obtained by burning sulphur or iron pyrites, the latter being thereby
converted into oxide of iron and sulphurous anhydride. In contact with
the oxygen of the atmosphere this gas does not pass into the higher
degree of oxidation, sulphuric anhydride, SO_{3}, and if it does form
sulphuric acid with water and the oxygen of the atmosphere, SO_{2} +
H_{2}O + O = H_{2}SO_{4}, it does so very slowly. With nitric acid (and
especially with nitrous acid, but not with nitrogen peroxide) and water,
sulphurous anhydride, on the contrary, very easily forms sulphuric acid,
and especially so when slightly heated (about 40°), the nitric acid (or,
better still, nitrous acid) being converted into nitric oxide--
3SO_{2} + 2NHO_{3} + 2H_{2}O = 2H_{2}SO_{4} + 2NO.
[63] The oxides of nitrogen naturally do not proceed directly from
oxygen and nitrogen by contact alone, because their formation
is accompanied by the absorption of a large quantity of heat,
for (_see_ Note 29) about 21,500 heat units are absorbed when 16
parts of oxygen and 14 parts of nitrogen combine; consequently
the decomposition of nitric oxide into oxygen and nitrogen
is accompanied by the evolution of this amount of heat; and
therefore with nitric oxide, as with all explosive substances and
mixtures, the reaction once started is able to proceed by itself.
In fact, Berthelot remarked the decomposition of nitric oxide
in the explosion of fulminate of mercury. This decomposition
does not take place spontaneously; substances even burn with
difficulty in nitric oxide, probably because a certain portion of
the nitric oxide in decomposing gives oxygen, which combines with
another portion of nitric oxide, and forms nitric peroxide, a
somewhat more stable compound of nitrogen and oxygen. The further
combinations of nitric oxide with oxygen all proceed with the
evolution of heat, and take place spontaneously by contact with
air alone. It is evident from these examples that the application
of thermochemical data is limited.
The presence of water is absolutely indispensable here, otherwise
sulphuric anhydride is formed, which combines with the oxides of
nitrogen (nitrous anhydride), forming a crystalline substance containing
oxides of nitrogen (_chamber crystals_, which will be described in