THE COMPOUNDS OF NITROGEN WITH HYDROGEN AND OXYGEN
[Illustration: FIG. 42.--The dry distillation of bones on a large scale.
The bones are heated in the vertical cylinders C (about 1-1/2 metre
high and 30 centimetres in diameter). The products of distillation pass
through the tubes T, into the condenser B, and receiver F. When the
distillation is completed the trap H is opened, and the burnt bones are
loaded into trucks V. The roof M is then opened, and the cylinders are
charged with a fresh quantity of bones. The ammonia water is preserved,
and goes to the preparation of ammoniacal salts, as described in the
following drawing.]
In the last chapter we saw that nitrogen does not directly combine
with hydrogen, but that a mixture of these gases in the presence of
hydrochloric acid gas, HCl, forms ammonium chloride, NH_{4}Cl, on the
passage of a series of electric sparks.[1] In ammonium chloride, HCl
is combined with NH_{3}, consequently N with H_{3} forms ammonia.[2]
Almost all the _nitrogenous substances of plants and animals_ evolve
ammonia when heated with an alkali. But even without the presence of
an alkali the majority of nitrogenous substances, when decomposed or
heated with a limited supply of air, evolve their nitrogen, if not
entirely, at all events partially, in the form of ammonia. When animal
substances such as skins, bones, flesh, hair, horns, &c., are heated
without access of air in iron retorts--they undergo what is termed
dry distillation. A portion of the resultant substances remains in the
retort and forms a carbonaceous residue, whilst the other portion, in
virtue of its volatility, escapes through the tube leading from the
retort. The vapours given off, on cooling, form a liquid which separates
into two layers; the one, which is oily, is composed of the so-called
animal oils (_oleum animale_): the other, an aqueous layer, contains
a solution of ammonia salts. If this solution be mixed with lime and
heated, the lime takes up the elements of carbonic acid from the ammonia
salts, and ammonia is evolved as a gas.[3] In ancient times ammonia
compounds were imported into Europe from Egypt, where they were prepared
from the soot obtained in the employment of camels' dung as fuel in
the locality of the temple of Jupiter Ammon (in Lybia), and therefore
the salt obtained was called 'sal-ammoniacale,' from which the name of
ammonia is derived. At the present time ammonia is obtained exclusively,
on a large scale, either from the products of the dry distillation of
animal or vegetable refuse, from urine, or from the ammoniacal liquors
collected in the destructive distillation of coal for the preparation
of coal gas. This ammoniacal liquor is placed in a retort with lime and
heated; the ammonia is then evolved together with steam.[4] In the arts,
only a small amount of ammonia is used in a free state--that is, in an
aqueous solution; the greater portion of it is converted into different
salts having technical uses, especially sal-ammoniac, NH_{4}Cl, and
ammonium sulphate, (NH_{4})_{2}SO_{4}. They are saline substances
which are formed because ammonia, NH_{3}, combines with all acids, HX,
forming ammonia salts, NH_{4}X. Sal-ammoniac, NH_{4}Cl, is a compound of
ammonia with hydrochloric acid. It is prepared by passing the vapours of
ammonia and water, evolved, as above described, from ammoniacal liquor,
into an aqueous solution of hydrochloric acid, and on evaporating the
solution sal-ammoniac is obtained in the form of soluble crystals[5]
resembling common salt in appearance and properties. Ammonia may be
very easily prepared _from_ this _sal-ammoniac_, NH_{4}Cl, as from any
other ammoniacal salt, by heating it with lime. Calcium hydroxide,
CaH_{2}O_{2}, as an alkali takes up the acid and sets free the
ammonia, forming calcium chloride, according to the equation 2NH_{4}Cl
+ CaH_{2}O_{2} = 2H_{2}O + CaCl_{2} + 2NH_{3}. In this reaction the
ammonia is evolved as a gas.[6]
[Illustration: FIG. 43.--Method of abstracting ammonia, on a large
scale, from ammonia water obtained at gas works by the dry distillation
of coal, or by the fermentation of urine, &c. This water is mixed with
lime and poured into the boiler C´´, and from thence into C´ and C
consecutively. The last boiler is heated directly over a furnace, and
hence no ammonia remains in solution after the liquid has been boiled
in it. The liquid is therefore then thrown away. The ammonia vapour and
steam pass from the boiler C, through the tube T, into the boiler C´,
and then into C´´, so that the solution in C´ becomes stronger than that
in C, and still stronger in C´´. The boilers are furnished with stirrers
A, A´, and A´´ to prevent the lime settling. From C´´ the ammonia and
steam pass through the tube T´´ into worm condensers surrounded with
cold water, thence into the Woulfe's bottle P, where the solution of
ammonia is collected, and finally the still uncondensed ammonia vapour
is led into the flat vessel R, containing acid which absorbs the last
traces of ammonia.]
[1] The ammonia in the air, water, and soil proceeds from the
decomposition of the nitrogenous substances of plants and animals,
and also probably from the reduction of nitrates. Ammonia is
always formed in the rusting of iron. Its formation in this case
depends in all probability on the decomposition of water, and on
the action of the hydrogen at the moment of its evolution on the
nitric acid contained in the air (Cloez), or on the formation of
ammonium nitrite, which takes place under many circumstances. The
evolution of vapours of ammonia compounds is sometimes observed
in the vicinity of volcanoes. At a red heat nitrogen combines
directly with B Ca Mg, and with many other metals, and these
compounds, when heated with a caustic alkali, or in the presence
of water, give ammonia (_see_ Chapter XIV., Note 14, and Chapter
XVII., Note 12). These are examples of the indirect combination of
nitrogen with hydrogen.
[2] If a silent discharge or a series of electric sparks be passed
through ammonia gas, it is decomposed into nitrogen and
hydrogen. This is a phenomenon of dissociation; therefore, a
series of sparks do not totally decompose the ammonia, but
leave a certain portion undecomposed. One volume of nitrogen
and three volumes of hydrogen are obtained from two volumes of
ammonia decomposed. Ramsay and Young (1884) investigated the
decomposition of NH_{3} under the action of heat, and showed
that at 500°, 1-1/2 p.c. is decomposed, at 600° about 18 p.c.,
at 800° 65 p.c., but these results were hardly free from the
influence of 'contact.' The _presence_ of free ammonia--that is,
ammonia not combined with acids--in a gas or aqueous solution
may be recognised by its characteristic smell. But many ammonia
salts do not possess this smell. However, on the addition of
an alkali (for instance, caustic lime, potash, or soda), they
evolve ammonia gas, especially when heated. The presence of
ammonia may be made visible by introducing a substance moistened
with strong hydrochloric acid into its neighbourhood. A white
cloud, or visible white vapour, then makes its appearance. This
depends on the fact that both ammonia and hydrochloric acid are
volatile, and on coming into contact with each other produce
solid sal-ammoniac, NH_{4}Cl, which forms a cloud. This test is
usually made by dipping a glass rod into hydrochloric acid, and
holding it over the vessel from which the ammonia is evolved. With
small amounts of ammonia this test is, however, untrustworthy,
as the white vapour is scarcely observable. In this case it is
best to take paper moistened with mercurous nitrate, HgNO_{3}.
This paper turns black in the presence of ammonia, owing to the
formation of a black compound of ammonia with mercurous oxide. The
smallest traces of ammonia (for instance, in river water) may be
detected by means of the so-called Nessler's reagent, containing
a solution of mercuric chloride and potassium iodide, which forms
a brown coloration or precipitate with the smallest quantities
of ammonia. It will be useful here to give the thermochemical
data (in thousands of units of heat, according to Thomsen), or
the quantities of heat _evolved_ in the formation of ammonia and
its compounds in quantities expressed by their formulæ. Thus, for
instance, (N + H_{3}) 26·7 indicates that 14 grams of nitrogen
in combining with 3 grams of hydrogen develop sufficient heat to
raise the temperature of 26·7 kilograms of water 1°. (NH_{3} +
nH_{2}O) 8·4 (heat of solution); (NH_{3},nH_{2}O + HCl,nH_{2}O)
12·3; (N + H_{4} + Cl) 90·6; (NH_{3} + HCl) 41·9.
[3] The same ammonia water is obtained, although in smaller quantities,
in the dry distillation of plants and of coal, which consists
of the remains of fossil plants. In all these cases the ammonia
proceeds from the destruction of the complex nitrogenous
substances occurring in plants and animals. The ammonia salts
employed in the arts are prepared by this method.
[4] The technical methods for the preparation of ammonia water, and for
the extraction of ammonia from it, are to a certain extent
explained in the figures accompanying the text.
[5] Usually these crystals are sublimed by heating them in crucibles or
pots, when the vapours of sal-ammoniac condense on the cold covers
as a crust, in which form the salt comes into the market.
[6] On a small scale ammonia may be prepared in a glass flask by mixing
equal parts by weight of slaked lime and finely-powdered
sal-ammoniac, the neck of the flask being connected with an
arrangement for drying the gas obtained. In this instance
neither calcium chloride nor sulphuric acid can be used for
drying the gas, since both these substances absorb ammonia, and
therefore solid caustic potash, which is capable of retaining
the water, is employed. The gas-conducting tube leading from the
desiccating apparatus is introduced into a mercury bath, if dry
gaseous ammonia be required, because water cannot be employed
in collecting ammonia gas. Ammonia was first obtained in this
dry state by Priestley, and its composition was investigated by
Berthollet at the end of the last century. Oxide of lead mixed
with sal-ammoniac (Isambert) evolves ammonia with still greater
ease than lime. The cause and process of the decomposition are
almost the same, 2PbO + 2NH_{4}Cl = Pb_{2}OCl_{2} + H_{2}O +
2NH_{3}. Lead oxychloride is (probably) formed.
It must be observed that all the complex nitrogenous substances of
plants, animals, and soils are decomposed when heated with an excess
of sulphuric acid, the whole of their nitrogen being converted into
ammonium sulphate, from which it may be liberated by treatment with
an excess of alkali. This reaction is so complete that it forms the
basis of Kjeldahl's method for estimating the amount of nitrogen in its
compounds.
Ammonia is a colourless gas, resembling those with which we are already
acquainted in its outward appearance, but clearly distinguishable from
any other gas by its very characteristic and pungent smell. It irritates
the eyes, and it is positively impossible to inhale it. Animals die in
it. Its density, referred to hydrogen, is 8·5; hence it is lighter than
air. It belongs to the class of gases which are easily liquefied.[7]
Faraday employed the following method for liquefying ammonia. Ammonia
when passed over dry silver chloride, AgCl, is absorbed by it to a
considerable extent, especially at low temperatures.[8] The solid
compound AgCl,3NH_{3} thus obtained is introduced into a bent tube
(fig. 45), whose open end c is then fused up. The compound is then
slightly heated at _a_, and the ammonia comes off, owing to the easy
dissociation of the compound. The other end of the tube is immersed in
a freezing mixture. The pressure of the gas coming off, combined with
the low temperature at one end of the tube, causes the ammonia evolved
to condense into a liquid, in which form it collects at the cold end of
the tube. If the heating be stopped, the silver chloride again absorbs
the ammonia. In this manner one tube may serve for repeated experiments.
Ammonia may also be liquefied by the ordinary methods--that is, by means
of pumping dry ammonia gas into a refrigerated space. Liquefied ammonia
is a colourless and very mobile liquid,[9] whose specific gravity at 0°
is 0·63 (E. Andréeff). At the temperature (about -70°) given by a mixture
of liquid carbonic anhydride and ether, liquid ammonia crystallises, and
in this form its odour is feeble, because at so low a temperature its
vapour tension is very inconsiderable. The boiling point (at a pressure
of 760 mm.) of liquid ammonia is about -32°. Hence this temperature may
be obtained at the ordinary pressure by the evaporation of liquefied
ammonia.
[Illustration: FIG. 45.--The liquefaction of ammonia in a thick bent
glass tube. A compound of chloride of silver and ammonia is placed in
the end _a_, and the end _c_ is then sealed up.]
[7] [Illustration: FIG. 44.--Carré's apparatus. Described in text.]
This is evident from the fact that its absolute boiling point lies
at about +130° (Chapter II., Note 29). It may therefore be
liquefied by pressure alone at the ordinary, and even at much
higher temperatures. The latent heat of evaporation of 17 parts
by weight of ammonia equals 4,400 units of heat, and hence
liquid ammonia may be employed for the production of cold.
Strong aqueous solutions of ammonia, which in parting with their
ammonia act in a similar manner, are not unfrequently employed
for this purpose. Suppose a saturated solution of ammonia to be
contained in a closed vessel furnished with a receiver. If the
ammoniacal solution be heated, the ammonia, with a small quantity
of water, will pass off from the solution, and in accumulating
in the apparatus will produce a considerable pressure, and will
therefore liquefy in the cooler portions of the receiver. Hence
liquid ammonia will be obtained in the receiver. The heating of
the vessel containing the aqueous solution of ammonia is then
stopped. After having been heated it contains only water, or a
solution poor in ammonia. When once it begins to cool the ammonia
vapours commence dissolving in it, the space becomes rarefied, and
a rapid vaporisation of the liquefied ammonia left in the receiver
takes place. In evaporating in the receiver it will cause the
temperature in it to fall considerably, and will itself pass into
the aqueous solution. In the end, the same ammoniacal solution as
originally taken is re-obtained. Thus, in this case, on heating
the vessel the pressure increases by itself, and on cooling it
diminishes, so that here heat directly replaces mechanical work.
This is the principle of the simplest forms of _Carré's ice-making
machines_, shown in fig. 44. C is a vessel made of boiler plates
into which the saturated solution of ammonia is poured; m is a
tube conducting the ammonia vapour to the receiver A. All parts
of the apparatus should be hermetically joined together, and
should be able to withstand a pressure reaching ten atmospheres.
The apparatus should be freed from air, which would otherwise
hinder the liquefaction of the ammonia. The process is carried on
as follows:--The apparatus is first so inclined that any liquid
remaining in A may flow into C. The vessel C is then placed
upon a stove F, and heated until the thermometer _t_ indicates
a temperature of 130° C. During this time the ammonia has been
expelled from C, and has liquefied in A. In order to facilitate
the liquefaction, the receiver A should be immersed in a tank of
water R (_see_ the left-hand drawing in fig. 44). After about
half an hour, when it may be supposed that the ammonia has been
expelled, the fire is removed from under C, and this is now
immersed in the tank of water R. The apparatus is represented in
this position in the right-hand drawing of fig. 44. The liquefied
ammonia then evaporates, and passes over into the water in C. This
causes the temperature of A to fall considerably. The substance
to be refrigerated is placed in a vessel G, in the cylindrical
space inside the receiver A. The refrigeration is also kept on for
about half an hour, and with an apparatus of ordinary dimensions
(containing about two litres of ammonia solution), five kilograms
of ice are produced by the consumption of one kilogram of coal. In
industrial works more complicated types of Carré's machines are
employed.
[8] Below 15° (according to Isambert), the compound AgCl,3NH_{3} is
formed, and above 20° the compound 2AgCl,3NH_{3}. The tension of
the ammonia evolved from the latter substance is equal to the
atmospheric pressure at 68°, whilst for AgCl,3NH_{3} the pressures
are equal at about 20°; consequently, at higher temperatures
it is greater than the atmospheric pressure, whilst at lower
temperatures the ammonia is absorbed and forms this compound.
Consequently, all the phenomena of dissociation are here clearly
to be observed. Joannis and Croisier (1894) investigated similar
compounds with AgBr, AgI, AgCN and AgNO_{3}, and found that
they all give definite compounds with NH_{3}, for instance
AgBr,3NH_{3}, 2AgBr,3NH_{3} and AgBr,2NH_{3}; they are all
colourless, solid substances which decompose under the atmospheric
pressure at +3·5, +34° and +51°.
[9] The liquefaction of ammonia may be accomplished without an increase
of pressure, by means of refrigeration alone, in a carefully
prepared mixture of ice and calcium chloride (because the absolute
boiling point of NH_{3} is high, about +130°). It may even take
place in the severe frosts of a Russian winter. The application of
liquid ammonia as a motive power for engines forms a problem which
has to a certain extent been solved by the French engineer Tellier.
Ammonia, containing, as it does, much hydrogen, is _capable of
combustion_; it does not, however, burn steadily, and sometimes not
at all, in ordinary atmospheric air. In pure oxygen it burns with a
greenish-yellow flame,[10] forming water, whilst the nitrogen set
free gives its oxygen compounds--that is, oxides of nitrogen. The
decomposition of ammonia into hydrogen and nitrogen not only takes
place at a red heat and under the action of electric sparks, but also
by means of many oxidising substances; for instance, by passing ammonia
through a tube containing red-hot copper oxide. The water thus formed
may be collected by substances absorbing it, and the quantity of
nitrogen may be measured in a gaseous form, and thus the composition
of ammonia determined. In this manner it is very easy to prove that
ammonia contains 3 parts by weight of hydrogen to 14 parts by weight of
nitrogen; and, by volume, 3 vols. of hydrogen and 1 vol. of nitrogen
form 2 vols. of ammonia.[11]
[10] The combustion of ammonia in oxygen may be effected by the aid of
platinum. A small quantity of an aqueous solution of ammonia,
containing about 20 p.c. of the gas, is poured into a wide-necked
beaker of about one litre capacity. A gas-conducting tube about
10 mm. in diameter, and supplying oxygen, is immersed in the
aqueous solution of ammonia. But before introducing the gas an
incandescent platinum spiral is placed in the beaker; the ammonia
in the presence of the platinum is oxidised and burns, whilst the
platinum wire becomes still more incandescent. The solution of
ammonia is heated, and oxygen passed through the solution. The
oxygen, as it bubbles off from the ammonia solution, carries with
it a part of the ammonia, and this mixture explodes on coming
into contact with the incandescent platinum. This is followed
by a certain cooling effect, owing to the combustion ceasing,
but after a short interval this is renewed, so that one feeble
explosion follows after another. During the period of oxidation
without explosion, white vapours of ammonium nitrite and
red-brown vapours of oxides of nitrogen make their appearance,
while during the explosion there is complete combustion and
consequently water and nitrogen are formed.
[11] This may be verified by their densities. Nitrogen is 14 times
denser than hydrogen, and ammonia is 8-1/2 times. If 3 volumes
of hydrogen with 1 volume of nitrogen gave 4 volumes of ammonia,
then these 4 volumes would weigh 17 times as much as 1 volume of
hydrogen; consequently 1 volume of ammonia would be 4-1/4 times
heavier than the same volume of hydrogen. But if these 4 volumes
only give 2 volumes of ammonia, the latter will be 8-1/2 times as
dense as hydrogen, which is found to be actually the case.
Ammonia is capable of combining with a number of substances, forming,
like water, substances of various degrees of stability. It is more
soluble than any of the gases yet described, both in water and in many
aqueous solutions. We have already seen, in the first chapter, that
one volume of water, at the ordinary temperature, dissolves about 700
vols. of ammonia gas. The great solubility of ammonia enables it to
be always kept ready for use in the form of an aqueous solution,[12]
which is commercially known as _spirits of hartshorn_. Ammonia water is
continually evolving ammoniacal vapour, and so has the characteristic
smell of ammonia itself. It is a very characteristic and important
fact that ammonia has an alkaline reaction, and colours litmus paper
blue, just like caustic potash or lime; it is therefore sometimes
called _caustic ammonia_ (volatile alkali). Acids may be saturated by
ammonia water or gas in exactly the same way as by any other alkali.
In this process _ammonia combines directly with acids_, and this forms
the most essential chemical reaction of this substance. If sulphuric,
nitric, acetic, or any other acid be brought into contact with ammonia
it absorbs it, and in so doing evolves a large amount of heat and forms
a compound having all the properties of a salt. Thus, for example,
sulphuric acid, H_{2}SO_{4}, in absorbing ammonia, forms (on evaporating
the solution) two salts, according to the relative quantities of ammonia
and acid. One salt is formed from NH_{3} + H_{2}SO_{4}, and consequently
has the composition NH_{5}SO_{4}, and the other is formed from 2NH_{3}
+ H_{2}SO_{4}, and its composition is therefore N_{2}H_{8}SO_{4}. The
former has an acid reaction and the latter a neutral reaction, and
they are called respectively acid ammonium sulphate (ammonium hydrogen
sulphate), and normal ammonium sulphate, or simply ammonium sulphate.
The same takes place in the action of all other acids; but certain of
them are able to form normal ammonium salts only, whilst others give
both acid and normal ammonium salts. This depends on the nature of
the acid and not on the ammonia, as we shall afterwards see. Ammonium
salts are very similar in appearance and in many of their properties
to metallic salts; for instance, sodium chloride, or table salt,
resembles sal-ammoniac, or ammonium chloride, not only in its outward
appearance but even in crystalline form, in its property of giving
precipitates with silver salts, in its solubility in water, and in
its evolving hydrochloric acid when heated with sulphuric acid--in a
word, a most perfect analogy is to be remarked in an entire series of
reactions. An analogy in composition is seen if sal-ammoniac, NH_{4}Cl,
be compared with table salt, NaCl; and the ammonium hydrogen sulphate,
NH_{4}HSO_{4}, with the sodium hydrogen sulphate, NaHSO_{4}; or ammonium
nitrate, NH_{4}NO_{3}, with sodium nitrate, NaNO_{3}.[13] It is seen,
on comparing the above compounds, that the part which sodium takes in
the sodium salts is played in ammonium salts by a group NH_{4}, which
is called _ammonium_. If table salt be called 'sodium chloride,' then
sal-ammoniac should be and is called 'ammonium chloride.'
[12] Aqueous solutions of ammonia are lighter than water, and at 15°,
taking water at 4° = 10,000, their specific gravity, as dependent
on _p_, or the percentage amount (by weight) of ammonia, is given
by the expression _s_ = 9,992-42·5_p_ + 0·21_p_^2; for instance,
with 10 p.c. _s_ = 9,587. If _t_ represents the temperature
between the limits of +10° and +20°, then the expression
(15-_t_)(1·5 + 0·14_p_) must be added to the formula for the
specific gravity. Solutions containing more than 24 p.c. have
not been sufficiently investigated in respect to the variation
of their specific gravity. It is, however, easy to obtain
more concentrated solutions, and at 0° solutions approaching
NH_{3},H_{2}O (48·6 p.c. NH_{3}) in their composition, and of
sp. gr. 0·85, may be prepared. But such solutions give up the
bulk of their ammonia at the ordinary temperature, so that more
than 24 p.c. NH_{3} is rarely contained in solution. Ammoniacal
solutions containing a considerable amount of ammonia give
ice-like crystals which seem to contain ammonia at temperatures
far below 0° (for instance, an 8 p.c. solution at -14°, the
strongest solutions at -48°). The whole of the ammonia may be
expelled from a solution by heating, even at a comparatively
low temperature; hence on heating aqueous solutions containing
ammonia a very strong solution of ammonia is obtained in the
distillate. Alcohol, ether, and many other liquids are also
capable of dissolving ammonia. Solutions of ammonia, when exposed
to the atmosphere, give off a part of their ammonia in accordance
with the laws of the solution of gases in liquids, which we have
already considered. But the ammoniacal solutions at the same time
absorb carbonic anhydride from the air, and ammonium carbonate
remains in the solution.
[Illustration: FIG. 46.--Apparatus for preparing solutions of
ammonia.]
Solutions of ammonia are required both for laboratory and factory
operations, and have therefore to be frequently prepared. For
this purpose the arrangement shown in fig. 46 is employed in the
laboratory. In works the same arrangement is used, only on a
larger scale (with earthenware or metallic vessels). The gas is
prepared in the retort, from whence it is led into the two-necked
globe A, and then through a series of Woulfe's bottles, B, C,
D, E. The impurities spurting over collect in A, and the gas is
dissolved in B, but the solution soon becomes saturated, and a
purer (washed) ammonia passes over into the following vessels,
in which only a pure solution is obtained. The bent funnel tube
in the retort preserves the apparatus from the possibility both
of the pressure of the gas evolved in it becoming too great
(when the gas escapes through it into the air), and also from
the pressure incidentally falling too low (for instance, owing
to a cooling effect, or from the reaction stopping). If this
takes place, the air passes into the retort, otherwise the liquid
from B would be drawn into A. The safety tubes in each Woulfe's
bottle, open at both ends, and immersed in the liquid, serve for
the same purpose. Without them, in case of an accidental stoppage
in the evolution of so soluble a gas as ammonia, the solution
would be sucked from one vessel to another--for instance, from
E into D, &c. In order to clearly see the necessity for _safety
tubes_ in a gas apparatus, it must be remembered that the
_gaseous pressure_ in the interior of the arrangement must exceed
the atmospheric pressure by the height of the sum of the columns
of liquid through which the gas has to pass.
[13] The analogy between the ammonium and sodium salts might seem to
be destroyed by the fact that the latter are formed from the
alkali or oxide and an acid, with the separation of water, whilst
the ammonium salts are directly formed from ammonia and an acid,
without the separation of water; but the analogy is restored
if we compare soda to ammonia water, and liken caustic soda to
a compound of ammonia with water. Then the very preparation of
ammonium salts from such a hydrate of ammonia will completely
resemble the preparation of sodium salts from soda. We may cite
as an example the action of hydrochloric acid on both substances.
NaHO + HCl = H_{2}O + NaCl
Sodium hydroxide Hydrochloric acid Water Table salt
NH_{4}HO + HCl = H_{2}0 + NH_{4}Cl
Ammonium hydroxide Hydrochloric acid Water Sal-ammoniac
Just as in soda the hydroxyl or aqueous radicle OH is replaced by
chlorine, so it is in ammonia hydrate.
The hypothesis that ammoniacal salts correspond with a complex metal
ammonium bears the name of the _ammonium theory_. It was enunciated
by the famous Swedish chemist Berzelius after the proposition made by
Ampère. The analogy admitted between ammonium and metals is probable,
owing to the fact that mercury is able to form an amalgam with ammonium
similar to that which it forms with sodium or many other metals. The
only difference between ammonium amalgam and sodium amalgam consists in
the instability of the ammonium, which easily decomposes into ammonia
and hydrogen.[14] Ammonium amalgam may be prepared from sodium amalgam.
If the latter be shaken up with a strong solution of sal-ammoniac, the
mercury swells up violently and loses its mobility whilst preserving its
metallic appearance. In so doing, the mercury dissolves ammonium--that
is, the sodium in the mercury is replaced by the ammonium, and replaces
it in the sal-ammoniac, forming sodium chloride, NH_{4}Cl + HgNa =
NaCl + HgNH_{4}. Naturally, the formation of ammonium amalgam does not
entirely prove the existence of ammonium itself in a separate state; but
it shows the possibility of this substance existing, and its analogy
with the metals, because only metals dissolve in mercury.[15] Ammonium
amalgam crystallises in cubes, three times heavier than water; it is
only stable in the cold, and particularly at very low temperatures. It
begins to decompose at the ordinary temperature, evolving ammonia and
hydrogen in the proportion of two volumes of ammonia and one volume
of hydrogen, NH_{4} = NH_{3} + H. By the action of water, ammonium
amalgam gives hydrogen and ammonia water, just as sodium amalgam gives
hydrogen and sodium hydroxide; and therefore, in accordance with
the ammonium theory, ammonia water must be looked on as containing
ammonium hydroxide, NH_{4}OH,[16] just as an aqueous solution of sodium
hydroxide, contains NaOH. The ammonium hydroxide, like ammonium itself,
is an unstable substance, which easily dissociates, and can only exist
in a free state at low temperatures.[17] Ordinary solutions of ammonia
must be looked on as the products of the dissociation of this hydroxide,
inasmuch as NH_{4}OH = NH_{3} + H_{2}O.
[14] Weyl (1864) by subjecting sodium to the action of ammonia at
the ordinary temperature and under considerable pressures,
obtained a liquid, which was subsequently investigated by Joannis
(1889), who confirmed the results obtained by Weyl. At 0° and the
atmospheric pressure the composition of this substance is Na +
5·3NH_{3}. The removal (at 0°) of ammonia from the liquid gives
a solid copper-red body having the composition NH_{3}Na. The
determination of the molecular weight of this substance by the
fall of the tension of liquid ammonia gave N_{2}H_{6}Na_{2}. It
is, therefore, free ammonium in which one H is replaced by Na.
The compound with potassium, obtained under the same conditions,
proved to have an analogous composition. By the decomposition of
NH_{3}Na at the ordinary temperature, Joannis (1891) obtained
hydrogen and sodium-amide NH_{2}Na in small colourless crystals
which were soluble in water. The addition of liquid ammonia to
metallic sodium and a saturated solution of sodium chloride,
gives NH_{2}Na_{2}Cl, and this substance is sal-ammoniac, in
which H_{2} is replaced by Na_{2}.
If pure oxygen be passed through a solution of these compounds
in ammonia at a temperature of about -50°, it is seen that the
gas is rapidly absorbed. The liquid gradually loses its dark
red colour and becomes lighter, and when it has become quite
colourless a gelatinous precipitate is thrown down. After the
removal of the ammonia, this precipitate dissolves easily in
water with a considerable evolution of heat, but without giving
off any gaseous products. The composition of the sodium compound
thus obtained is NH_{2}Na_{2}HO, which shows that it is a hydrate
of bisodium-ammonium. Thus, although free ammonium has not been
obtained, still a sodium substitution product of it is known
which corresponds to it as a salt to a hydrate. Ammonium amalgam
was originally obtained in exactly the same way as sodium amalgam
(Davy); namely, a piece of sal-ammoniac was taken, and moistened
with water (in order to render it a conductor of electricity). A
cavity was made in it, into which mercury was poured, and it was
laid on a sheet of platinum connected with the positive pole of a
galvanic battery, while the negative pole was put into connection
with the mercury. On passing a current the mercury increased
considerably in volume, and became plastic, whilst preserving
its metallic appearance, just as would be the case were the
sal-ammoniac replaced by a lump of a sodium salt or of many other
metals. In the analogous decomposition of common metallic salts,
the metal contained in a given salt separates out at the negative
pole, immersed in mercury, by which the metal is dissolved. A
similar phenomenon is observed in the case of sal-ammoniac; the
elements of ammonium, NH_{4}, in this case are also collected in
the mercury, and are retained by it for a certain time.
[15] We may mention, however, that under particular conditions hydrogen
is also capable of forming an amalgam resembling the amalgam of
ammonium. If an amalgam of zinc be shaken up with an aqueous
solution of platinum chloride, without access of air, then a
spongy mass is formed which easily decomposes, with the evolution
of hydrogen.
[16] We saw above that the solubility of ammonia in water at low
temperatures attains to the molecular ratio NH_{3} + H_{2}O, in
which these substances are contained in caustic ammonia, and
perhaps it may be possible at exceedingly low temperatures to
obtain ammonium hydroxide, NH_{4}HO, in a solid form. Regarding
solutions as dissociated definite compounds, we should see a
confirmation of this view in the property shown by ammonia of
being extremely soluble in water, and in so doing of approaching
to the limit NH_{4}HO.
[17] In confirmation of the truth of this conclusion we may cite the
remarkable fact that there exist, in a free state and as
comparatively stable compounds, a series of alkaline hydroxides,
NR_{4}HO, which are perfectly analogous to ammonium hydroxide,
and present a striking resemblance to it and to sodium hydroxide,
with the only difference that the hydrogen in NH_{4}HO is
replaced by complex groups, R = CH_{3}, C_{2}H_{5}, &c., for
instance N(CH_{3})_{4}HO. Details will be found in organic
chemistry.
All ammoniacal salts _decompose at a red heat_ into ammonia and an acid,
which, on cooling in contact with each other, re-combine together. If
the acid be non-volatile, the ammoniacal salt, when heated, evolves the
ammonia, leaving the non-volatile acid behind; if the acid be volatile,
then, on heating, both the acid and ammonia volatilise together, and
on cooling re-combine into the salt which originally served for the
formation of their vapours.[18]
[18] The fact that ammoniacal salts are decomposed when ignited, and not
simply sublimed, may be proved by a direct experiment with
sal-ammoniac, NH_{4}Cl, which in a state of vapour is decomposed
into ammonia, NH_{3}, and hydrochloric acid, HCl, as will be
explained in the following chapter. The readiness with which
ammonium salts decompose is seen from the fact that a solution
of ammonium oxalate is decomposed with the evolution of ammonia
even at -1°. Dilute solutions of ammonium salts, when boiled give
aqueous vapour having an alkaline reaction, owing to the presence
of free ammonia given off from the salt.
Ammonia is not only capable of combining with acids, but also with many
salts, as was seen from its forming definite compounds, AgCl,3NH_{3}
and 2AgCl,3NH_{3}, with silver chloride. Just as ammonia is absorbed
by various oxygen salts of the metals, so also is it absorbed by the
chlorine, iodine, and bromine compounds of many metals, and in so doing
evolves heat. Certain of these compounds part with their ammonia even
when left exposed to the air, but others only do so at a red heat; many
give up their ammonia when dissolved, whilst others dissolve without
decomposition, and when evaporated separate from their solutions
unchanged. All these facts only indicate that ammoniacal, like aqueous,
compounds dissociate with greater or lesser facility.[19] Certain
metallic oxides also absorb ammonia and are dissolved in ammonia water.
Such are, for instance, the oxides of zinc, nickel, copper, and many
others; the majority of such compounds are unstable. The property of
ammonia of combining with certain oxides explains its action on certain
metals.[20] By reason of such action, copper vessels are not suitable
for holding liquids containing ammonia. Iron is not acted on by such
liquids.
[19] Isambert studied the dissociation of ammoniacal compounds, as we
have seen in Note 8, and showed that at low temperatures many
salts are able to combine with a still greater amount of
ammonia, which proves an entire analogy with hydrates; and as
in this case it is easy to isolate the definite compounds, and
as the least possible tension of ammonia is greater than that
of water, therefore the ammoniacal compounds present a great
and peculiar interest, as a means for explaining the nature of
aqueous solutions and as a confirmation of the hypothesis of
the formation of definite compounds in them; for these reasons
we shall frequently refer to these compounds in the further
exposition of this work.
[20] Chapter V., Note 2.
The similarity between the relation of ammonia and water to salts and
other substances is more especially marked in those cases in which
the salt is capable of combining with both ammonia and water. Take,
for example, copper sulphate, CuSO_{4}. As we saw in Chapter I., it
gives with water blue crystals, CuSO_{4},5H_{2}O; but it also absorbs
ammonia in the same molecular proportion, forming a blue substance,
CuSO_{4},5NH_{3}, and therefore the ammonia combining with salts may be
termed _ammonia of crystallisation_.
Such are the _reactions of combination_ proper to ammonia. Let us
now turn our attention to the reactions of substitution proper to
this substance. If ammonia be passed through a heated tube containing
metallic sodium, hydrogen is evolved, and a compound is obtained
containing ammonia in which one atom of hydrogen is replaced by an
atom of sodium, NH_{2}Na (according to the equation NH_{3} + Na =
NH_{2}Na + H). This body is termed sodium amide. We shall afterwards
see that iodine and chlorine are also capable of directly displacing
hydrogen from ammonia, and of replacing it. In fact, the hydrogen of
ammonia may be replaced in many ways by different elements. If in this
replacement NH_{2} remains, the resultant substances NH_{2}R are called
_amides_, whilst the substitution products, NHR_{2}, in which only NH
remains, are called _imides_,[20 bis] and those in which none of the
ammoniacal hydrogen remains, NR_{3}, are known as _nitrides_. Free
amidogen, N_{2}H_{4}, is now known in a state of hydration under the
name of hydrazine;[21] it combines with acids and resembles ammonia in
this respect. In the action of different substances on ammonia it is
the _hydrogen that is substituted_, whilst the nitrogen remains in the
resultant compound, so to say, untouched. The same phenomenon is to be
observed in the action of various substances on water. In the majority
of cases the reactions of water consist in the hydrogen being evolved,
and in its being replaced by different elements. This also takes place,
as we have seen, in acids in which the hydrogen is easily displaced by
metals. This chemical mobility of hydrogen is perhaps connected with the
great lightness of the atoms of this element.
[20 bis] Imide, NH, has not been obtained in a free state, but its
hydrochloric acid salt, NHHCl, has apparently been obtained (1890)
by Maumené by igniting the double bichloride of platinum and
ammonium chloride, PtCl_{2}NH_{4}Cl = Pt + 2HCl + NHHCl. It is
soluble in water, and crystallises from its solution in hexagonal
rhombic prisms. It gives a double salt with FeCl_{3} of the
composition FeCl_{3}3NHHCl. The salt NHHCl is similar (isomeric)
with the first possible product of the metalepsis of ammonia,
NH_{2}Cl, although it does not resemble it in any of its
properties.
[21] Free _amidogen_ or _hydrazine_, N_{2}H_{4}, or 2NH_{2}, was
prepared by Curtius (1887) by means of ethyl diazoacetate,
or triazoacetic acid. Curtius and Jay (1889) showed that
triazoacetic acid, CHN_{2}.COOH (the formula should be tripled),
when heated with water or a mineral acid, gives (quantitatively)
oxalic acid and amidogen (hydrazine), CHN_{2}.COOH + 2H_{2}O =
C_{2}O_{2}(OH)_{2} + N_{2}H_{4}--_i.e._ (empirically), the oxygen
of the water replaces the nitrogen of the azoacetic acid. The
amidogen is thus obtained in the form of a salt. With acids,
amidogen forms very stable salts of the two types N_{2}H_{4}HX
and N_[2]H_{4}H_{2}X_{2}, as, for example, with HCl, H_{2}SO_{4},
&c. These salts are easily crystallised; in acid solutions
they act as powerful reducing agents, evolving nitrogen; when
ignited they are decomposed into ammoniacal salts, nitrogen,
and hydrogen; with nitrites they evolve nitrogen. The sulphate
N_{2}H_{4},H_{2}SO_{4} is sparingly soluble in cold water (3
parts in 100 of water), but is very soluble in hot water; its
specific gravity is 1·378, it fuses at 254° with decomposition.
The hydrochloride N_{2}H_{4},2HCl crystallises in octahedra, is
very soluble in water, but not in alcohol; it fuses at 198°,
evolving hydrogen chloride and forming the salt N_{2}H_{4}HCl;
when rapidly heated it decomposes with an explosion; with
platinic chloride it immediately evolves nitrogen, forming
platinous chloride. By the action of alkalis the salts
N_{2}H_{4},2HX give _hydrate of amidogen_, N_{2}H_{4},H_{2}O,
which is a fuming liquid (specific gravity 1·03), boiling at
119°, almost without odour, and whose aqueous solution corrodes
glass and india-rubber, has an alkaline taste and poisonous
properties. The reducing capacities of the hydrate are clearly
seen from the fact that it reduces the metals platinum and
silver from their solutions. With mercuric oxide it explodes. It
reacts directly with the aldehydes RO, forming N_{2}R_{2} and
water; for example, with benzaldehydes it gives the very stable
insoluble _benzalazine_ (C_{6}H_{5}CHN)_{2} of a yellow colour.
We may add that hydrazine often forms double salts; for example,
MgSO_{4}N_{2}H_{4}H_{2}SO_{4} or KClN_{2}H_{4}HCl, and that it is
also formed by the action of nitrous acid upon aldehyde-ammonia.
The products of the substitution of the hydrogen in hydrazine by
hydrocarbon groups R (R = CH_{3}, C_{2}H_{5}, C_{6}H_{5}, &c.)
were obtained before hydrazine itself; for example, NHRNH_{2},
NR_{2}NH_{2}, and (NRH)_{2}.
The heat of solution of the sulphuric acid salt (1 part in 200
and 300 parts of water at 10°·8) is equal to -8·7 C. According
to Berthelot and Matigon (1892), the heat of neutralisation of
hydrazine by sulphuric acid is +5·5 C and by hydrochloric acid
+5·2 C. Thus hydrazine is a very feeble base, for its heat of
saturation is not only lower than that of ammonia (+12·4 C.
for HCl), but even below that of hydroxylamine (+9·3 C.) The
heat of formation from the elements of hydrated hydrazine -9·5 C
was deduced from the heat of combustion, determined by burning
N_{2}H_{4}H_{2}SO_{4} in a calorimetric bomb, +127·7 C. Thus
hydrazine is an endothermal compound; its passage into ammonia
by the combination of hydrogen is accompanied by the evolution
of 51·5 C. In the presence of an acid these figures were greater
by +14·4 C. Hence the direct converse passage from ammonia into
hydrazine is impossible. As regards the passage of hydroxylamine
into hydrazine, it would be accompanied by the evolution of heat
(+21·5 C.) in an aqueous solution.
Amidogen must be regarded as a compound which stands to ammonia
in the same relation as hydrogen peroxide stands to water.
Water, H(OH), gives, according to the law of substitution,
as was clearly to be expected, (OH)(OH)--that is, peroxide
of hydrogen is the free radicle of water (hydroxyl). So also
ammonia, H(NH_{2}), forms hydrazine, (NH_{2})(NH_{2})--that is,
the free radicle of ammonia, NH_{2}, or amidogen. In the case of
phosphorus a similar substance, as we shall afterwards see, has
long been known under the name of liquid phosphuretted hydrogen,
P_{2}H_{4}.
In practical chemistry[21 bis] ammonia is often employed, not only
for saturating acids, but also for effecting reactions of double
decomposition with salts, and especially for separating insoluble
basic hydroxides from soluble salts. Let MHO stand for an insoluble
basic hydroxide and HX for an acid. The salt formed by them will have a
composition MHO + XH-H_{2}O = MX. If aqueous ammonia, NH_{4}OH, be added
to a solution of this salt, the ammonia will change places with the
metal M, and thus form the insoluble basic hydroxide, or, as it is said,
give a precipitate.
MX + NH_{4}(OH) = NH_{4}X + MHO
Salt of the metal. Aqueous ammonia. Ammonium salt. Basic hydrate.
In solution In solution In solution As precipitate
[21 bis] In practice, the applications of ammonia are very varied. The
use of ammonia as a stimulant, in the forms of the so-called
'smelling salts' or of spirits of hartshorn, in cases of
faintness, &c., is known to everyone. The volatile carbonate of
ammonium, or a mixture of an ammonium salt with an alkali, is also
employed for this purpose. Ammonia also produces a well-known
stimulating effect when rubbed on the skin, for which reason it is
sometimes employed for external applications. Thus, for instance,
the well-known volatile salve is prepared from any liquid oil
shaken up with a solution of ammonia. A portion of the oil is thus
transformed into a soapy substance. The solubility of greasy
substances in ammonia, which proceeds from the formation both of
emulsions and soaps, explains its use in extracting grease spots.
It is also employed as an external application for stings from
insects, and for bites from poisonous snakes, and in general in
medicine. It is also remarkable that in cases of drunkenness a few
drops of ammonia in water taken internally rapidly renders a
person sober. A large quantity of ammonia is used in dyeing,
either for the solution of certain dyes--for example, carmine--or
for changing the tints of others, or else for neutralising the
action of acids. It is also employed in the manufacture of
artificial pearls. For this purpose the small scales of a peculiar
small fish are mixed with ammonia, and the liquid so obtained is
blown into small hollow glass beads shaped like pearls.
In nature and the arts, however, ammonium salts, and not free
ammonia, are most frequently employed. In this form a portion of
that _nitrogen_ which is necessary for the formation of albuminous
substances is _supplied to plants_. Owing to this, a large
quantity of ammonium sulphate is now employed as a fertilising
substance. But the same effect may be produced by nitre, or by
animal refuse, which in decomposing gives ammonia. For this
reason, an ammoniacal (hydrogen) compound may be introduced into
the soil in the spring which will be converted into a nitrate
(oxygen salt) in the summer.
Thus, for instance, if aqueous ammonia is added to a solution of a salt
of aluminium, then alumina hydrate is separated out as a colourless
gelatinous precipitate.[22]
[22] As certain basic hydrates form peculiar compounds with ammonia,
in some cases it happens that the first portions of ammonia
added to a solution of a salt produce a precipitate, whilst
the addition of a fresh quantity of ammonia dissolves this
precipitate if the ammoniacal compound of the base be soluble in
water. This, for example, takes place with the copper salts. But
alumina does not dissolve under these circumstances.
In order to grasp the relation between ammonia and the oxygen
compounds of nitrogen it is necessary to recognise the general
_law of substitution_, applicable to all cases of substitution
between elements,[23] and therefore showing what may be the cases of
substitution between oxygen and hydrogen as component parts of water.
The law of substitution may be deduced from mechanical principles if
the molecule be conceived as a system of elementary atoms occurring
in a certain chemical and mechanical equilibrium. By likening the
molecule to a system of bodies in a state of motion--for instance,
to the sum total of the sun, planets, and satellites, existing in
conditions of mobile equilibrium--then we should expect the action
of one part, in this system, to be equal and opposite to the other,
according to Newton's third law of mechanics. Hence, given a molecule
of a compound, for instance, H_{2}O, NH_{3}, NaCl, HCl, &c., its every
two parts must in a chemical sense represent two things somewhat
alike in force and properties, and therefore _every two parts into
which a molecule of a compound may be divided are capable of replacing
each other_. In order that the application of the law should become
clear it is evident that among compounds the most stable should be
chosen. We will therefore take hydrochloric acid and water as the
most stable compounds of hydrogen.[24] According to the above law of
substitution, if the elements H and Cl are able to form a molecule, HCl,
and a stable one, they are able to replace each other. And, indeed,
we shall afterwards see (Chapter XI.) that in a number of instances a
substitution between hydrogen and chlorine can take place. Given RH,
then RCl is possible, because HCl exists and is stable. The molecule of
water, H_{2}O, may be divided in two ways, because it contains 3 atoms:
into H and (HO) on the one hand, and into H_{2} and O on the other.
Consequently, being given RH, its substitution products will be R(HO)
according to the first form, and R_{2}O according to the second; being
given RH_{2}, its corresponding substitution products will be RH(OH),
R(OH)_{2}, RO, (RH)_{2}O, &c. The group (OH) is the same hydroxyl or
aqueous radicle which we have already mentioned in the third chapter
as a component part of hydroxides and alkalis--for instance, Na(OH),
Ca(OH)_{2}, &c. It is evident, judging from H(HO) and HCl, that (OH)
can be substituted by Cl, because both are replaceable by H; and this
is of common occurrence in chemistry, because metallic chlorides--for
example, NaCl and NH_{4}Cl--correspond with hydroxides of the alkalis
Na(OH) or NH_{4}(OH). In hydrocarbons--for instance, C_{2}H_{6}--the
hydrogen is replaceable by chlorine and by hydroxyl. Thus ordinary
alcohol is C_{2}H_{6}, in which one atom of H is replaced by (OH); that
is, C_{2}H_{5}(OH). It is evident that the replacement of hydrogen by
hydroxyl essentially forms the phenomenon of oxidation, because RH gives
R(OH), or RHO. Hydrogen peroxide may in this sense be regarded as water
in which the hydrogen is replaced by hydroxyl; H(OH) gives (OH)_{2} or
H_{2}O_{2}. The other form of substitution--namely, that of O in the
place of H_{2}--is also a common chemical phenomenon. Thus alcohol,
C_{2}H_{6}O, or C_{2}H_{5}(OH), when oxidising in the air, gives acetic
acid, C_{2}H_{4}O_{2}, or C_{2}H_{3}O(OH), in which H_{2} is replaced
by O.
[23] When the element chlorine, as we shall afterwards more fully
learn, replaces the element hydrogen, the reaction by which
such an exchange is accomplished proceeds as a substitution, AH
+ Cl_{2} = ACl + HCl, so that two substances, AH and chlorine,
react on each other, and two substances, ACl and HCl, are
formed; and further, two molecules react on each other, and
two others are formed. The reaction proceeds very easily, but
the substitution of one element, _A_, by another, _X_, does
not always proceed with such ease, clearness, or simplicity.
The substitution between oxygen and hydrogen is very rarely
accomplished by the reaction of the free elements, but the
substitution between these elements, one for another, forms
the most common case of oxidation and reduction. In speaking
of the law of substitution, I have in view the substitution of
the elements one by another, and not the direct reaction of
substitution. The law of substitution determines the cycle of the
combinations of a given element, if a few of its compounds (for
instance, the hydrogen compounds) be known. A development of the
conceptions of the law of substitution may be found in my lecture
given at the Royal Institution in London, 1889.
[24] If hydrogen peroxide be taken as a starting point, then still
higher forms of oxidation than those corresponding with water
should be looked for. They should possess the properties of
hydrogen peroxide, especially that of parting with their oxygen
with extreme ease (even by contact). Such compounds are known.
Pernitric, persulphuric, and similar acids present these
properties, as we shall see in describing them.
In the further course of this work we shall have occasion to refer to
the law of substitution for explaining many chemical phenomena and
relations.
We will now apply these conceptions to ammonia in order to see its
relation to the oxygen compounds of nitrogen. It is evident that many
substances should be obtainable from ammonia, NH_{3}, or aqueous
ammonia, NH_{4}(OH), by substituting their hydrogen by hydroxyl,
or H_{2} by oxygen. And such is the case. The two extreme cases of
such substitution will be as follows: (1) One atom of H in NH_{3} is
substituted by (OH), and NH_{2}(OH) is produced. Such a substance, still
containing much hydrogen, should have many of the properties of ammonia.
It is known under the name of _hydroxylamine_,[25] and, in fact, is
capable, like ammonia, of giving salts with acids; for example, with
hydrochloric acid, NH_{3}(OH)Cl--which is a substance corresponding to
sal-ammoniac, in which one atom of hydrogen is replaced by hydroxyl.[25
bis] (2) The other extreme case of substitution is that given by
ammonium hydroxide, NH_{4}(OH), when the whole of the hydrogen of the
ammonium is replaced by oxygen; and, as ammonium contains 4 atoms of
hydrogen, the highest oxygen compound should be NO_{2}(OH), or NHO_{3},
as we find to be really the case, for NHO_{3} is nitric acid, exhibiting
the highest degree of oxidation of nitrogen.[26] If instead of the two
extreme aspects of substitution we take an intermediate one, we obtain
the intermediate oxygen compounds of nitrogen. For instance, N(OH)_{3}
is orthonitrous acid,[27] to which corresponds nitrous acid, NO(OH), or
NHO_{2}, equal to N(OH)_{3}-H_{2}O, and nitrous anhydride, N_{2}O_{3} =
2N(OH)_{3}-3H_{2}O. Thus nitrogen gives a series of oxygen compounds,
which we will proceed to describe. We will, however, first show by
two examples that in the first place the passage of ammonia into the
oxygen compounds of nitrogen up to nitric acid, as well as the converse
preparation of ammonia (and consequently of the intermediate compounds
also) from nitric acid, are reactions which proceed directly and easily
under many circumstances, and in the second place that the above general
principle of substitution gives the possibility of understanding many,
at first sight unexpected and complex, relations and transformations,
such as the preparation of hydronitrous acid, HN_{3}. In nature the
matter is complicated by a number of influences and circumstances, but
in the law the relations are presented in their simplest aspect.
[25] The compound of hydroxylamine with hydrochloric acid has the
composition NH_{2}(OH)HCl = NH_{4}ClO--that is, it is as it were
oxidised sal-ammoniac. It was prepared by Lossen in 1865 by the
action of tin and hydrochloric acid in the presence of water on
a substance called ethyl nitrate, in which case the hydrogen
liberated from the hydrochloric acid by the tin acts upon the
elements of nitric acid--
C_{2}H_{5}·NO_{3} + 6H + HCl = NH_{4}OCl + H_{2}O + C_{2}H_{5}·OH
Ethyl nitrate Hydrogen Hydroxylamine Water Alcohol
from + HCl
HCl and Sn
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.