Chapter VI: The Compounds of Nitrogen with Hydrogen and Oxygen (3)

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