Chapter VI: The Compounds of Nitrogen with Hydrogen and Oxygen (1)

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]

[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.

[7]

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.

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