Chapter VI: The Compounds of Nitrogen with Hydrogen and Oxygen (2)

Thus in this case the nitric acid is deoxidised, not directly
into nitrogen, but into hydroxylamine. Hydroxylamine is also
formed by passing nitric oxide, NO, into a mixture of tin and
hydrochloric acid--that is, by the action of the hydrogen evolved
on the nitric oxide, NO + 3H + HCl = NH_{4}OCl--and in many other
cases. According to Lossen's method, a mixture of 30 parts of
ethyl nitrate, 120 parts of tin, and 40 parts of a solution of
hydrochloric acid of sp. gr. 1·06 are taken. After a certain
time the reaction commences spontaneously. When the reaction
has ceased the tin is separated by means of hydrogen sulphide,
the solution is evaporated, and a large amount of sal-ammoniac
is thus obtained (owing to the further action of hydrogen on
the hydroxylamine compound, the hydrogen taking up oxygen from
it and forming water); a solution ultimately remains containing
the hydroxylamine salt; this salt is dissolved in anhydrous
alcohol and purified by the addition of platinum chloride, which
precipitates any ammonium salt still remaining in the solution.
After concentrating the alcoholic solution the hydroxylamine
hydrochloride separates in crystals. This substance melts at
about 150°, and in so doing decomposes into nitrogen, hydrogen
chloride, water, and sal-ammoniac. A sulphuric acid compound of
hydroxylamine may be obtained by mixing a solution of the above
salt with sulphuric acid. The sulphate is also soluble in water
like the hydrochloride; this shows that hydroxylamine, like
ammonia itself, forms a series of salts in which one acid may
be substituted for another. It might he expected that by mixing
a strong solution of a hydroxylamine salt with a solution of a
caustic alkali hydroxylamine itself would be liberated, just
as an ammonia salt under these circumstances evolves ammonia;
but the liberated hydroxylamine is immediately decomposed with
the formation of nitrogen and ammonia (and probably nitrous
oxide), 3NH_{3}O = NH_{3} + 3H_{2}O + N_{2}. Dilute solutions
give the same reaction, although very slowly, but by decomposing
a solution of the sulphate with barium hydroxide a certain
amount of hydroxylamine is obtained in solution (it is partly
decomposed). Hydroxylamine in aqueous solution, like ammonia,
precipitates basic hydrates, and it deoxidises the oxides of
copper, silver, and other metals. Free hydroxylamine was obtained
by Lobry de Bruyn (1891). It is a solid, colourless, crystalline
substance, without odour, which does not melt below 27°. It has
the property of dissolving metallic salts; for instance, sodium
chloride. Hydroxylamine, when rapidly heated with platinum,
decomposes with a flash and the formation of a yellow flame.
It is almost insoluble in ordinary solvents like chloroform,
benzine, acetic ether, and carbon bisulphide. Its aqueous
solutions are tolerably stable, contain up to 60 per cent.
(sp. gr. 1·15 at 20°), and may be kept for many weeks without
undergoing any change. Lobry de Bruyn used the hydrochloric
salt to prepare pure hydroxylamine. The salt was first treated
with sodium methylate (CH_{3}NaO), and then methyl alcohol was
added to the mixture. The precipitated sodium chloride was
separated from the solution by filtration. (The methyl alcohol
is added to prevent the precipitated chloride of sodium from
coating the insoluble hydrochloric salt of hydroxylamine.) The
methyl alcohol was driven off under a pressure 150-200 mm., and
after extracting a further portion of methyl alcohol by ether
and several fractional distillations, a solution was obtained
containing 70 per cent. of free hydroxylamine, 8 per cent.
water, 9·9 per cent. chloride of sodium, and 12·1 per cent. of
the hydrochloric salt of hydroxylamine. Pure free hydroxylamine,
NH_{3}O, is obtained by distilling under a pressure of 60 mm.;
it then boils at 70°, and solidifies in a condenser cooled to
0° in the form of long needles. It melts at 33°, boils at 58°
under a pressure of 22 mm., and has a sp. gr. of about 1·235
(Brühl). Under the action of NaHO it gives NH_{3} and NHO_{2}
or N_{2}O, and forms nitric acid (Kolotoff, 1893) under the
action of oxidising agents. Hydroxylamine is obtained in a great
number of cases, for instance by the action of tin on dilute
nitric acid, and also by the action of zinc on ethyl nitrate and
dilute hydrochloric acid, &c. The relation between hydroxylamine,
NH_{2}(OH), and nitrous acid, NO(OH), which is so clear in the
sense of the law of substitutions, becomes a reality in those
cases when reducing agents act on salts of nitrous acid. Thus
Raschig (1888) proposed the following method for the preparation
of the hydroxylamine sulphate. A mixture of strong solutions of
potassium nitrite, KNO_{2}, and hydroxide, KHO, in molecular
proportions, is prepared and cooled. An excess of sulphurous
anhydride is then passed into the mixture, and the solution
boiled for a long time. A mixture of the sulphates of potassium
and hydroxylamine is thus obtained: KNO_{2} + KHO + 2SO_{2} +
2H_{2}O = NH_{2}(OH),H_{2}SO_{4} + K_{2}SO_{4}. The salts may be
separated from each other by crystallisation.

[25 bis] In order to illustrate the application of the law of
substitution to a given case, and to show the connection between
ammonia and the oxides of nitrogen, let us consider the possible
products of an oxygen and hydroxyl substitution in caustic
ammonia, NH_{4}(OH). It is evident that the substitution of H by
OH can give: (1) NH_{3}(OH)_{2}; (2) NH_{2}(OH)_{3}; (3)
NH(OH)_{4}; and (4) N(OH)_{5}. They should all, like caustic
ammonia itself, easily part with water and form products
(hydroxylic) of the oxidation of ammonia. The first of them is the
hydrate of hydroxylamine, NH_{2}(OH) + H_{2}O; the second,
NH(OH)_{2} + H_{2}O (and also the substance NH(OH)_{4} or
NH_{3}O_{2}), containing, as it does, both hydrogen and oxygen, is
able to part with all its hydrogen in the form of water (which
could not be done by the first product, since it contained too
little oxygen), forming, as the ultimate product,
2NH_{2}(OH)_{3}-5H_{2}O = N_{2}O--that is, it corresponds with
nitrous oxide, or the lower degree of the oxidation of nitrogen.
So, also, nitrous anhydride corresponds with the third of the
above products, 2NH(OH)_{4}-5H_{2}O = N_{2}O_{3}, and nitric
anhydride with the fourth, 2N(OH)_{5}-5H_{2}O = N_{2}O_{5}. As, in
these three equations, two molecules of the substitution products
(-5H_{2}O) are taken, it is also possible to combine two different
products in one equation. For instance, the third and fourth
products: NH(OH)_{4} + N(OH)_{5}-5H_{2}O corresponds to N_{2}O_{4}
or 2NO_{2}, that is, to peroxide of nitrogen. Thus all the five
(see later) oxides of nitrogen, N_{2}O, NO, N_{2}O_{3}, NO_{2},
and N_{2}O_{5}, may be deduced from ammonia. The above may be
expressed in a general form by the equation (it should be remarked
that the composition of all the substitution products of caustic
ammonia may be expressed by NH_{3}O{5-_a_}, where _a_ varies
between 0 and 4):

NH_{5}O_{5 - _a_} + NH_{5}O{5 - _b_} - 5H_{2}O
= N_{2}O_{5 - (_a_ + _b_)},

where _a_ + _b_ can evidently be not greater than 5; when _a_ +
_b_ = 5 we have N_{2}--nitrogen, when = 4 we have N_{2}O nitrous
oxide; when _a_ + _b_ = 3 we have N_{2}O_{2} or NO--nitric oxide,
and so on to N_{2}O_{5}, when _a_ + _b_ = 0. Besides which it is
evident that intermediate products may correspond with (and hence
also break up into) different starting points; for instance,
N_{2}O is obtained when _a_ + _b_ = 2, and this may occur either
when _a_ = 0 (nitric acid), and _b_ = 2 (hydroxylamine), or when
_a_ = _b_ = 1 (the third of the above substitution products).

[26] Nitric acid corresponds with the anhydride N_{2}O_{5}, which will
afterwards be described, but which must be regarded as the
highest saline oxide of nitrogen, just as Na_{2}O (and the
hydroxide NaHO) in the case of sodium, although sodium forms
a peroxide possessing the property of parting with its oxygen
with the same ease as hydrogen peroxide, if not on heating,
at all events in reactions--for instance, with acids. So also
nitric acid has its corresponding peroxide, which may be called
pernitric acid. Its composition is not well known--probably
NHO_{4}--so that its corresponding anhydride would be N_{2}O_{7}.
It is formed by the action of a silent discharge on a mixture
of nitrogen and oxygen, so that a portion of its oxygen is
in a state similar to that in ozone. The instability of this
substance (obtained by Hautefeuille, Chappuis, and Berthelot),
which easily splits up with the formation of nitric peroxide, and
its resemblance to persulphuric acid, which we shall afterwards
describe, will permit our passing over the consideration of the
little that is further known concerning it.

[27] Phosphorus (Chapter XIX.) gives the hydride PH_{3}, corresponding
with ammonia, NH_{3}, and forms phosphorous acid, PH_{3}O_{3},
which is analogous to nitrous acid, just as phosphoric acid is
to nitric acid; but phosphoric (or, better, orthophosphoric)
acid, PH_{3}O_{4}, is able to lose water and give pyro-and
meta-phosphoric acids. The latter is equal to the ortho-acid
minus water = PHO_{3}, and therefore nitric acid, NHO_{3},
is really meta-nitric acid. So also nitrous acid, HNO_{2},
is meta-nitrous (anhydrous) acid, and thus the ortho-acid is
NH_{3}O_{3} = N(OH)_{3}. Hence for nitric acid we should expect
to find, besides the ordinary or meta-nitric acid, HNO_{3}
(= 1/2N_{2}O_{3},H_{2}O), and ortho-nitric acid, H_{3}NO_{4}
(= 1/2N_{2}O_{3},3H_{2}O), an intermediate pyro-nitric acid,
N_{2}H_{4}O_{7}, corresponding to pyrophosphoric acid,
P_{2}H_{4}O_{7}. We shall see (for instance, in Chapter XVI.,
Note 21) that in nitric acid there is indeed an inclination
of the ordinary salts (of the meta-acid), MNO_{3}, to combine
with bases M_{2}O, and to approximate to the composition of
ortho-compounds which are equal to meta-compound and bases
(MNO_{3} + M_{2}O = M_{3}NO_{4}).

1. It is easy to prove the possibility of the oxidation of ammonia into nitric acid by passing a mixture of ammonia and air over heated spongy platinum. This causes the oxidation of the ammonia, nitric acid being formed, which partially combines with the excess of ammonia.

The converse passage of nitric acid into ammonia is effected by the action of hydrogen at the moment of its evolution.[28] Thus metallic aluminium, evolving hydrogen from a solution of caustic soda, is able to completely convert nitric acid added to the mixture (as a salt, because the alkali gives a salt with the nitric acid) into ammonia, NHO_{3} + 8H = NH_{3} + 3H_{2}O.

2. In 1890 Curtius in Germany obtained a gaseous substance of the composition HN_{3} (hydrogen trinitride), having the distinctive properties of an acid, and giving, like hydrochloric acid, salts; for example, a sodium salt, NaN_{3}; ammonium salt, NH_{4}N_{3} = N_{4}H_{4}; barium salt, Ba(N_{3})_{2}, &c., which he therefore named hydronitrous acid, _HN__{3}.[28 bis] The extraordinary composition of the compound (ammonia, NH_{3}, contains one N atom and three H atoms; in HN_{3}, on the contrary, there are three N atoms and one H atom), the facile decomposition of its salts with an explosion, and above all its distinctly acid character (an aqueous solution shows a strong acid reaction to litmus), not only indicated the importance of this unexpected discovery, but at first gave rise to some perplexity as to the nature of the substance obtained, for the relations in which HN_{3} stood to other simple compounds of nitrogen which had long been known was not at all evident, and the scientific spirit especially requires that there should be a distinct bond between every innovation, every fresh discovery, and that which is already firmly established and known, for upon this basis is founded that apparently paradoxical union in science of a conservative stability with an irresistible and never-ceasing improvement. This missing, connection between the newly discovered hydronitrous acid, HN_{3}, and the long known ammonia, NH_{3}, and nitric acid, HNO_{3}, may be found in the law of substitution, starting from the well-known properties and composition of nitric acid and ammonia, as I mentioned in the 'Journal of the Russian Physico-Chemical Society' (1890). The essence of the matter lies in the fact that to the hydrate of ammonium, or caustic ammonia, NH_{4}OH, there should correspond, according to the law of substitution, an ortho-nitric acid (_see_ Note 27), H_{3}NO_{4} = NO(OH)_{3}, which equals NH_{4}(OH) with the substitution in it of (_a_) two atoms of hydrogen by oxygen (O--H_{2}) and (_b_) two atoms of hydrogen by the aqueous radicle (OH--H). Ordinary or meta-nitric acid is merely this ortho-nitric acid minus water. To ortho-nitric acid there should correspond the ammoniacal salts: mono-substituted, H_{2}NH_{4}NO_{4}; bi-substituted, H(NH_{4})_{2}NO_{4}; and tri-substituted, (NH_{4})_{3}NO_{4}. These salts, containing as they do hydrogen and oxygen, like many similar ammoniacal salts (see, for instance, Chapter IX.--Cyanides), are able to part with them in the form of water. Then from the first salt we have H_{2}NH_{4}NO_{4}-4H_{2}O = N_{2}O--nitrous oxide, and from the second H(NH_{4})_{2}NO4-4H_{2}O = HN_{3}--hydronitrous acid, and from the third (NH_{4})_{3}NO-4H_{2}O = N_{4}H_{4}--the ammonium salt of the same acid. The composition of HN_{3} should be thus understood, whilst its acid properties are explained by the fact that the water (4H_{2}O) from H(NH_{4})_{2}N_O{4} is formed at the expense of the hydrogen of the ammonium and oxygen of the nitric acid, so that there remains the same hydrogen as in nitric acid, or that which may be replaced by metals and give salts. Moreover, nitrogen undoubtedly belongs to that category of metalloids which give acids, like chlorine and carbon, and therefore, under the influence of three of its atoms, one atom of hydrogen acquires those properties which it has in acids, just as in HCN (hydrocyanic acid) the hydrogen has received these properties under the influence of the carbon and nitrogen (and HN_{3} may be regarded as HCN where C has been replaced by N_{2}). Moreover, besides explaining the composition and acid properties of HN_{3}, the above method gives the possibility of foretelling the closeness of the bond between hydronitrous acid and nitrous oxide, for N_{2}O + NH_{3} = HN_{3} + H_{2}O. This reaction, which was foreseen from the above considerations, was accomplished by Wislicenus (1892) by the synthesis of the sodium salt, by taking the amide of sodium, NH_{2}Na (obtained by heating Na in a current of NH_{3}), and acting upon it (when heated) with nitrous oxide, N_{2}O, when 2NH_{2}Na + N_{2}O = NaN_{3} + NaHO + NH_{3}. The resultant salt, NaN_{3}, gives hydronitrous acid when acted upon by sulphuric acid, NaN_{3} + H_{2}SO_{4} = NaHSO_{4} +HN_{3}. The latter gives, with the corresponding solutions of their salts, the insoluble (and easily explosive) salts of silver, AgN_{3} (insoluble, like AgCl or AgCN), and lead, Pb(N_{3})_{2}.

[28] The formation of ammonia is observed in many cases of oxidation by
means of nitric acid. This substance is even formed in the action
of nitric acid on tin, especially if dilute acid be employed in
the cold. A still more considerable amount of ammonia is obtained
if, in the action of nitric acid, there are conditions directly
tending to the evolution of hydrogen, which then reduces the acid
to ammonia; for instance, in the action of zinc on a mixture of
nitric and sulphuric acids.

[28 bis] Curtius started with benzoylhydrazine, C_{6}H_{5}CONHNH_{2}
(hydrazine, see Note 20 bis). (This substance is obtained by
the action of hydrated hydrazine on the compound ether of benzoic
acid). Benzoylhydrazine under the action of nitrous acid gives
benzoylazoimide and water:

C_{6}H_{5}CONHNH_{2} + NO_{2}H = C_{6}H_{5}CON_{3} + 2H_{2}O.

Benzoylazoimide when treated with sodium alcoholate gives the
sodium salt of hydronitrous acid:

C_{6}H_{5}CON_{3} + C_{2}H_{3}ONa
= C_{6}H_{5}O_{2}C_{2}H_{3} + NaN_{3}.

The addition of ether to the resultant solution precipitates the
NaN_{3}, and this salt when treated with sulphuric acid gives
gaseous hydronitrous acid, HN_{3}. It has an acrid smell, and is
easily soluble in water. The aqueous solution exhibits a strongly
acid reaction. Metals dissolve in this solution and give the
corresponding salts. With hydronitrous acid gaseous ammonia forms
a white cloud, consisting of the salt of ammonium, NH_{4}N_{3}.
This salt separates out from an alcoholic solution in the form of
white lustrous scales. The salts of hydronitrous acid are obtained
by a reaction of substitution with the sodium or ammonium salts.
In this manner Curtius obtained and studied the salts of silver
(AgN_{3}), mercury (HgN_{3}), lead (PbN_{6}), barium (BaN_{6}).
With hydrazine, N_{2}H_{4}, hydronitrous acid forms saline
compounds in the composition of which there are one or two
particles of N_{3}H per one particle of hydrazine; thus N_{5}H_{5}
and N_{8}H_{6}. The first was obtained in an almost pure form. It
crystallises from an aqueous solution in dense, volatile, lustrous
prisms (up to 1 in. long), which fuse at 50°, and deliquesce in
the air; from a solution in boiling alcohol it separates out in
bright crystalline plates. This salt, N_{5}H_{5}, has the same
empirical composition, NH, as the ammonium salt of hydronitrous
acid, N_{4}H_{4}, and imide; but their molecules and structure are
different. Curtius also obtained (1893) hydronitrous acid by
passing the vapour of N_{2}O_{5} (evolved by the action of HNO_{3}
on As_{2}O_{3}) into a solution of hydrazine, N_{2}H_{4}.
Similarly Angeli, by acting upon a saturated solution of silver
nitrite with a strong solution of hydrazine, obtained the
explosive AgN_{3} in the form of a precipitate, and this reaction,
which is based upon the equation N_{2}H_{4} + NHO_{2} = HN_{3} +
2H_{2}O, proceeds so easily that it forms an experiment for the
lecture table. A thermal investigation of hydronitrous acid by
Berthelot and Matignon gave the following figures for the heat of
solution of the ammonium salt N_{3}HNH_{3} (1 grm. in 100 parts of
water)-708 C., and for the heat of neutralisation by barium
hydrate +10·0 C., and by ammonia +8·2 C. The heat of combustion of
N_{4}H_{4} (+163·8 C. at a constant vol.) gives the heat of
formation of the salt N_{4}H_{4} (solid) as -25·3 C. and
(solution)-32·3 C.; this explains the explosive nature of this
compound. In its heat of formation from the elements N_{3}H =-62·6
C., this compound differs from all the hydrogen compounds of
nitrogen in having a maximum absorption of heat, which explains
its instability.

The compounds of nitrogen with oxygen present an excellent example of the law of multiple proportions, because they contain, for 14 parts by weight of nitrogen, 8, 16, 24, 32, and 40 parts respectively by weight of oxygen. The composition of these compounds is as follows:--

N_{2}O, nitrous oxide; hydrate NHO.
N_{2}O_{2}, nitric oxide, NO.
N_{2}O_{3}, nitrous anhydride; hydrate NHO_{2}.
N_{2}O_{4}, peroxide of nitrogen, NO_{2}.
N_{2}O_{5}, nitric anhydride; hydrate NHO_{3}.

Of these compounds,[29] nitrous and nitric oxides, peroxide of nitrogen, and nitric acid, NHO_{3}, are characterised as being the most stable. _The lower oxides, when coming into contact with the higher, may give the intermediate forms_; for instance, NO and NO_{2} form N_{2}O_{3}, _and the intermediate oxides may, in splitting up, give a higher and lower oxide_. So N_{2}O_{4} gives N_{2}O_{3} and N_{2}O_{5}, or, in the presence of water, their hydrates.

[29] According to the thermochemical determinations of Favre, Thomsen,
and more especially of Berthelot, it follows that, in the
formation of such quantities of the oxides of nitrogen as express
their formulæ, if gaseous nitrogen and oxygen be taken as the
starting points, and if the compounds formed be also gaseous, the
following amounts of heat, expressed in thousands of heat units,
are _absorbed_ (hence a minus sign):--

N_{2}O N_{2}O_{2} N_{2}O_{3} N_{2}O_{4} N_{2}O_{5}
-21 -43 -22 -5 -1
-22 +21 +17 +4

The difference is given in the lower line. For example, if
N_{2}, or 28 grams of nitrogen, combine with O--that is, with 16
grams of oxygen--then 21,000 units of heat are absorbed, that
is, sufficient heat to raise 21,000 grams of water through 1°.
Naturally, direct observations are impossible in this case; but
if charcoal, phosphorus, or similar substances are burnt both in
nitrous oxide and in oxygen, and the heat evolved is observed
in both cases, then the difference (more heat will be evolved
in burning in nitrous oxide) gives the figures required. If
N_{2}O_{2}, by combining with O_{2}, gives N_{2}O_{4}, then, as
is seen from the table, heat should be developed, namely, 38,000
units of heat, or NO + O = 19,000 units of heat. The differences
given in the table show that the maximum absorption of heat
corresponds with nitric oxide, and that the higher oxides are
formed from it with evolution of heat. If liquid nitric acid,
NHO_{3}, were decomposed into N + O_{3} + H, then 41,000 heat
units would be required; that is, an evolution of heat takes
place in its formation from the gases. It should be observed
that the formation of ammonia, NH_{3}, from the gases N + H_{3}
evolves 12·2 thousand heat units.

We have already seen that, under certain conditions, nitrogen combines 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}.

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}.

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]