Chapter V: Nitrogen and Air (1)

Gaseous _nitrogen_ forms about four-fifths (by volume) of the atmosphere; consequently the air contains an exceedingly large mass of it. Whilst entering in so considerable a quantity into the composition of air, nitrogen does not seem to play any active part in the atmosphere, the chemical action of which is mainly dependent on the oxygen it contains. But this is not an entirely correct idea, because animal life cannot exist in pure oxygen, in which animals pass into an abnormal state and die; and the nitrogen of the air, although slowly, forms diverse compounds, many of which play a most important part in nature, especially in the life of organisms. However, neither plants[1] nor animals directly absorb the nitrogen of the air, but take it up from already prepared nitrogenous compounds; further, plants are nourished by the nitrogenous substances contained in the soil and water, and animals by the nitrogenous substances contained in plants and in other animals. Atmospheric electricity is capable of aiding the passage of gaseous nitrogen into nitrogenous compounds, as we shall afterwards see, and the resultant substances are carried to the soil by rain, where they serve for the nourishment of plants. Plentiful harvests, fine crops of hay, vigorous growth of trees--other conditions being equal--are only obtained when the soil contains _ready prepared nitrogenous compounds_, consisting either of those which occur in air and water, or of the residues of the decomposition of other plants or animals (as in manure). The nitrogenous substances contained in animals have their origin in those substances which are formed in plants. Thus the nitrogen of the atmosphere is the origin of all the nitrogenous substances occurring in animals and plants, although not directly so, but after first combining with the other elements of air.

[1] See Note 15 bis.

The nitrogenous compounds which enter into the composition of plants and animals are of primary importance; no vegetable or animal cell--that is, the elementary form of organism--exists without containing a nitrogenous substance, and moreover organic life manifests itself primarily in these nitrogenous substances. The germs, seeds, and those parts by which cells multiply themselves abound in nitrogenous substances; the sum total of the phenomena which are proper to organisms depend primarily on the chemical properties of the nitrogenous substances which enter into their composition. It will be sufficient, for instance, to point out the fact that vegetable and animal organisms, clearly distinguishable as such, are characterised by a different degree of energy in their nature, and at the same time by a difference in the amount of nitrogenous substances they contain. In plants, which compared with animals possess but little activity, being incapable of independent movement, &c., the amount of nitrogen is very much less than in animals, whose tissues are almost exclusively formed of nitrogenous substances. It is remarkable that the nitrogenous parts of plants, chiefly of the lower orders, sometimes present both forms and properties which approach to those of animal organisms; for example, the zoospores of sea-weeds, or those parts by means of which the latter multiply themselves. These zoospores on leaving the sea-weed in many respects resemble the lower orders of animal life, having, like the latter, the property of moving. They also approach the animal kingdom in their composition, their outer coating containing nitrogenous matter. Directly the zoospore becomes covered with that non-nitrogenous or cellular coating which is proper to all the ordinary cells of plants, it loses all resemblance to an animal organism and becomes a small plant. It may be thought from this that the cause of the difference in the vital processes of animals and plants is the different amount of nitrogenous substances they contain. The nitrogenous substances which occur in plants and animals appertain to a series of exceedingly complex and very changeable chemical compounds; their elementary composition alone shows this; besides nitrogen, they contain carbon, hydrogen, oxygen, and sulphur. Being distinguished by a very great instability under many conditions in which other compounds remain unchanged, these substances are fitted for those perpetual changes which form the first condition of vital activity. These complex and changeable nitrogenous substances of the organism are called _proteïd substances_. The white of eggs is a familiar example of such a substance. They are also contained in the flesh of animals, the curdy elements of milk, the glutinous matter of wheaten flour, or so-called gluten, which forms the chief component of macaroni, &c.

Nitrogen occurs in the earth's crust, in compounds either forming the remains of plants and animals, or derived from the nitrogen of the atmosphere as a consequence of its combination with the other component parts of the air. It is not found in other forms in the earth's crust; so that nitrogen must be considered, in contradistinction to oxygen, as an element which is purely superficial, and does not extend to the depths of the earth.[1 bis]

[1 bis] The reason why there are no other nitrogenous substances within
the earth's mass beyond those which have come there with the
remains of organisms, and from the air with rain-water, must be
looked for in two circumstances. In the first place, in the
instability of many nitrogenous compounds, which are liable to
break up with the formation of gaseous nitrogen; and in the second
place in the fact that the salts of nitric acid, forming the
product of the action of air on many nitrogenous and especially
organic compounds, are very soluble in water, and on penetrating
into the depths of the earth (with water) give up their oxygen. The
result of the changes of the nitrogenous organic substances which
fall into the earth is without doubt frequently, if not invariably,
the formation of gaseous nitrogen. Thus the gas evolved from coal
always contains much nitrogen (together with marsh gas, carbonic
anhydride, and other gases).

_Nitrogen is liberated_ in a free state in the decomposition of the _nitrogenous organic substances_ entering into the composition of organisms--for instance, on their combustion. All organic substances burn when heated to redness with oxygen (or substances readily yielding it, such as oxide of copper); the oxygen combines with the carbon, sulphur, and hydrogen, and the nitrogen is evolved in a free state, because at a high temperature it does not form any stable compound, but remains uncombined. Carbonic anhydride and water are formed from the carbon and hydrogen respectively, and therefore to obtain pure nitrogen it is necessary to remove the carbonic anhydride from the gaseous products obtained. This may be done very easily by the action of alkalis--for instance, caustic soda. The amount of nitrogen in organic substances is determined by a method founded on this.

It is also very easy to obtain _nitrogen from air_, because oxygen combines with many substances. Either phosphorus or metallic copper is usually employed for removing the oxygen from air, but, naturally, a number of other substances may also be used. If a small saucer on which a piece of phosphorus is laid be placed on a cork floating on water, and the phosphorus be lighted, and the whole covered with a glass bell jar, then the air under the jar will be deprived of its oxygen, and nitrogen only will remain, owing to which, on cooling, the water will rise to a certain extent in the bell jar. The same object (procuring nitrogen from air) is attained much more conveniently and perfectly by passing air through a red-hot tube containing copper filings. At a red heat, metallic copper combines with oxygen and gives a black powder of copper oxide. If the layer of copper be sufficiently long and the current of air slow, all the oxygen will be absorbed, and nitrogen alone will pass from the tube.[2]

[2] Copper (best as turnings, which present a large surface) absorbs
oxygen, forming CuO, at the ordinary temperature in the presence
of solutions of acids, or, better still, in the presence of a
solution of ammonia, when it forms a bluish-violet solution of
oxide of copper in ammonia. Nitrogen is very easily procured by
this method. A flask filled with copper turnings is closed with a
cork furnished with a funnel and stopcock. A solution of ammonia
is poured into the funnel, and caused to drop slowly upon the
copper. If at the same time a current of air be slowly passed
through the flask (from a gasholder), then all the oxygen will
be absorbed from it and the nitrogen will pass from the flask.
It should be washed with water to retain any ammonia that may be
carried off with it.

Nitrogen may also be procured from many of its _compounds with oxygen[3] and hydrogen_,[4] but the best fitted for this purpose is a saline mixture containing, on the one hand, a compound of nitrogen with oxygen, termed nitrous anhydride, N_{2}O_{3}, and on the other hand, ammonia, NH_{3}--that is, a compound of nitrogen with hydrogen. By heating such a mixture, the oxygen of the nitrous anhydride combines with the hydrogen of the ammonia, forming water, and gaseous nitrogen is evolved, 2NH_{3} + N_{2}O_{3} = 3H_{2}O + N_{4}. Nitrogen is procured by this method in the following manner:--A solution of caustic potash is saturated with nitrous anhydride, by which means potassium nitrite is formed. On the other hand, a solution of hydrochloric acid saturated with ammonia is prepared; a saline substance called sal-ammoniac, NH_{4}Cl, is thus formed in the solution. The two solutions thus prepared are mixed together and heated. Reaction takes place according to the equation KNO_{2} + NH_{4}Cl = KCl + 2H_{2}O + N_{2}. This reaction proceeds in virtue of the fact that potassium nitrite and ammonium chloride are salts which, on interchanging their metals, give potassium chloride and ammonium nitrite, NH_{4}NO_{2}, which breaks up into water and nitrogen. This reaction does not take place without the aid of heat, but it proceeds very easily at a moderate temperature. Of the resultant substances, the nitrogen only is gaseous. Pure nitrogen may be obtained by drying the resultant gas and passing it through a solution of sulphuric acid (to absorb a certain quantity of ammonia which is evolved in the reaction).[4 bis]

[3] The oxygen compounds of nitrogen (for example, N_{2}O, NO, NO_{2})
are decomposed at a red heat by themselves, and under the action
of red-hot copper, iron, sodium, &c., they give up their oxygen
to the metals, leaving the nitrogen free. According to Meyer and
Langer (1885), nitrous oxide, N_{2}O, decomposes below 900°,
although not completely.

[4] Chlorine and bromine (in excess), as well as bleaching powder
(hypochlorites), take up the hydrogen from ammonia, NH_{3},
leaving nitrogen. Nitrogen is best procured from ammonia by the
action of a solution of sodium hypobromite on solid sal-ammoniac.

[4 bis] Lord Rayleigh in 1894, when determining the weight of a volume
of carefully purified nitrogen by weighing it in one and the same
globe, found that the gas obtained from air, by the action of
incandescent copper (or iron or by removing the oxygen by ferrous
oxide) was always 1/200 heavier than the nitrogen obtained from its
compounds, for instance, from the oxide or suboxide of nitrogen,
decomposed by incandescent pulverulent iron or from the ammonia
salt of nitrous acid. For the nitrogen procured from air, he
obtained, at 0° and 760·4 mm. pressure, a weight = 2·310 grms.,
while for the nitrogen obtained from its compounds, 2·299 grms.
This difference of about 1/200 could not be explained by the
nitrogen not having been well purified, or by inaccuracy of
experiment, and was the means for the remarkable discovery of the
presence of a heavy gas in air, which will be mentioned in Note
16 bis.

Nitrogen is a gaseous substance which does not differ much in physical properties from air; its density, referred to hydrogen, is approximately equal to 14--that is, it is slightly lighter than air, its density referred to air being 0·972; one litre of nitrogen weighs 1·257 gram. Nitrogen mixed with oxygen, which is slightly heavier than air, forms air. It is a gas which, like oxygen and hydrogen, is liquefied with difficulty, and is but little soluble in water and other liquids. Its absolute boiling point[5] is about -140°; above this temperature it is not liquefiable by pressure, and at lower temperatures it remains a gas at a pressure of 50 atmospheres. Liquid nitrogen boils at -193°, so that it may be employed as a source of great cold. At about -203°, in vaporising under a decrease of pressure, nitrogen solidifies into a colourless snow-like mass. Nitrogen does not burn,[5 bis] does not support combustion, is not absorbed by any of the reagents used in gas analysis, at least at the ordinary temperature--in a word, it presents a whole series of negative chemical properties; this is expressed by saying that this element has no energy for combination. Although it is capable of forming compounds both with oxygen and hydrogen as well as with carbon, yet these compounds are only formed under particular circumstances, to which we will directly turn our attention. At a red heat nitrogen combines with boron, titanium, and silicon, barium, magnesium, &c., forming very stable nitrogenous compounds,[6] whose properties are entirely different from those of nitrogen with hydrogen, oxygen and carbon. However, the combination of nitrogen with carbon, although it does not take place directly between the elements at a red heat, yet proceeds with comparative ease by heating a mixture of charcoal with an alkaline carbonate, especially potassium carbonate or barium carbonate, to redness, carbo-nitrides or cyanides of the metals being formed; for instance, K_{2}CO_{3} + 4C + N_{2} = 2KCN + 3CO.[7]

[5] See Chapter II. Note 29.

[5 bis] See Note 11 bis.

[6] The combination of boron with nitrogen is accompanied by the
evolution of sufficient heat to raise the mass to redness;
titanium combines so easily with nitrogen that it is difficult
to obtain it free from that element; magnesium easily absorbs
nitrogen at a red heat. It is a remarkable and instructive fact
that these compounds of nitrogen are very stable and non-volatile.
Carbon (C = 12) with nitrogen gives cyanogen, C_{2}N_{2}, which
is gaseous and very unstable, and whose molecule is not large,
whilst boron (B = 11) forms a nitrogenous compound which is solid,
non-volatile, and very stable. Its composition, BN, is similar
to that of cyanogen, but its molecular weight, B_{n}N_{n}, is
probably greater. Its composition, like that of N_{2}Mg_{3},
NNa_{3}, N_{2}Hg_{3} and of many of the metallic nitrides,
corresponds to ammonia with the substitution of all its hydrogen
by a metal. In my opinion, a detailed study of the transformations
of the nitrides now known, should lead to the discovery of many
facts in the history of nitrogen.

[7] This reaction, so far as is known, does not proceed beyond a
certain limit, probably because cyanogen, CN, itself breaks up
into carbon and nitrogen.

Nitrogen is found with oxygen in the air, but they do not readily combine. Cavendish, however, in the last century, showed that _nitrogen combines with oxygen under the influence of a series of electric sparks_. Electric sparks in passing through a moist[8] mixture of nitrogen and oxygen cause these elements to combine, forming reddish-brown fumes of oxides of nitrogen,[9] which form nitric acid,[10] NHO_{3}. The presence of the latter is easily recognised, not only from its reddening litmus paper, but also from its acting as a powerful oxidiser even of mercury. Conditions similar to these occur in nature, during a thunderstorm or in other electrical discharges which take place in the atmosphere; whence it may be taken for granted that air and rain-water always contain traces of nitric and nitrous acids.[11] Besides which Crookes (1892) showed that under certain circumstances and when electricity of high potential[11 bis] passes through the air, the combination of nitrogen with oxygen is accompanied by the formation of a true flame. This was also observed previously (1880) during the passage of electrical discharges through the air.

[8] Frémy and Becquerel took dry air, and observed the formation of
brown vapours of oxides of nitrogen on the passage of sparks.

[9] If a mixture of one volume of nitrogen and fourteen volumes of
hydrogen be burnt, then water and a considerable quantity of
nitric acid are formed. It may be partly due to this that a
certain quantity of nitric acid is produced in the slow oxidation
of nitrogenous substances in an excess of air. This is especially
facilitated by the presence of an alkali with which the nitric
acid formed can combine. If a galvanic current be passed through
water containing the nitrogen and oxygen of the air in solution,
then the hydrogen and oxygen set free combine with the nitrogen,
forming ammonia and nitric acid.

When copper is oxidised at the expense of the air at the ordinary
temperature in the presence of ammonia, oxygen is absorbed, not
only for combination with the copper, but also for the formation
of nitric acid.

The combination of nitrogen with oxygen, even, for example, by
the action of electric sparks, is not accompanied by an explosion
or rapid combination, as in the action of a spark on a mixture
of oxygen and hydrogen. This is explained by the fact that heat
is not evolved in the combination of nitrogen with oxygen, but
is absorbed--an expenditure of energy is required, there is no
evolution of energy. In fact, there will not be the transmission
of heat from particle to particle which occurs in the explosion
of detonating gas. Each spark will aid the formation of a certain
quantity of the compound of oxygen and nitrogen, but will not
excite the same in the neighbouring particles. In other words, the
combination of hydrogen with oxygen is an exothermal reaction, and
the combination of nitrogen with oxygen an endothermal reaction.

A condition particularly favourable for the oxidation of nitrogen
is the explosion of detonating gas and air if the former be
_in excess_. If a mixture of two volumes of detonating gas and
one volume of air be exploded, then one-tenth of the air is
converted into nitric acid, and consequently after the explosion
has taken place there remain only nine-tenths of the volume of
air originally taken. If a large proportion of air be taken--for
instance, four volumes of air to two volumes of detonating
gas--then the temperature of the explosion is lowered, the
volume of air taken remains unchanged, and no nitric acid is
formed. This gives a rule to be observed in making use of the
eudiometer--namely that to weaken the force of the explosion
not less than an equal volume of air should be added to the
explosive mixture. On the other hand a large excess must not be
taken as no explosion would then ensue (_see_ Chapter III. Note
34). Probably in the future means will be found for obtaining
compounds of nitrogen on a large industrial scale by the aid of
electric discharges, and by making use of the inexhaustible mass
of nitrogen in the atmosphere.

[10] In reality nitric oxide, NO, is first formed, but with oxygen and
water it gives (brown fumes) nitrous anhydride, which, as we
shall afterwards learn, in the presence of water and oxygen gives
nitric acid.

[11] The nitric acid contained in the soil, river water (Chapter I.,
Note 2), wells, &c., proceeds (like carbonic anhydride) from the
oxidation of organic compounds which have fallen into water,
soil, &c.

[11 bis] Crookes employed a current of 15 ampères and 65 volts, and
passed it through an induction coil with 330 vibrations per
second, and obtained a flame between the poles placed at a
distance of 46 mm. which after the appearance of the arc and flame
could be increased to 200 mm. A platinum wire fused in the flame.

Further observations showed that under the influence of electrical discharges,[12] silent as well as with sparks, nitrogen is able to enter into many reactions with hydrogen and with many hydrocarbons; although these reactions cannot be effected by exposure to a red heat. Thus, for instance, a series of electric sparks passed through a mixture of nitrogen and hydrogen causes them to combine and _form ammonia_[13] or nitrogen hydride, NH_{3}, composed of one volume of nitrogen and three volumes of hydrogen. This combination is limited to the formation of 6 per cent. of ammonia, because ammonia is decomposed, although not entirely (94/100) by electric sparks. This signifies that under the influence of an electrical discharge the reaction NH_{3} = N + 3H is reversible, consequently it is a dissociation, and in it a state of equilibrium is arrived at. The equilibrium may be destroyed by the addition of gaseous hydrochloric acid, HCl, because with ammonia it forms a solid saline compound, sal-ammoniac, NH_{4}Cl, which (being formed from a gaseous mixture of 3H, N, and HCl) fixes the ammonia. The remaining mass of nitrogen and hydrogen, under the action of the sparks, again forms ammonia, and in this manner _solid sal-ammoniac is obtained to the end by the action of a series of electric sparks on a mixture of gaseous_ N, H_{3}, _and_ HCl.[14] Berthelot (1876) showed that under the action of a silent discharge many non-nitrogenous organic substances (benzene, C_{6}H_{6}, cellulose in the form of paper, resin, glucose, C_{6}H_{10}O_{5}, and others) absorb nitrogen and form complex nitrogenous compounds, which are capable, like albuminous substances, of evolving their nitrogen as ammonia when heated with alkalis.[15]

[12] This property of nitrogen, which under normal conditions is
inactive, leads to the idea that under the influence of an
electric discharge gaseous nitrogen changes in its properties;
if not permanently like oxygen (electrolysed oxygen or ozone
does not react on nitrogen, according to Berthelot), it may be
temporarily at the moment of the action of the discharge, just
as some substances under the action of heat are permanently
affected (that is, when once changed remain so--for instance,
white phosphorus passes into red, &c.), whilst others are only
temporarily altered (the dissociation of S_{6} into S_{2} or
of sal-ammoniac into ammonia and hydrochloric acid). Such a
proposition is favoured by the fact that nitrogen gives two kinds
of spectra, with which we shall afterwards become acquainted. It
may be that the molecules N_{2} then give less complex molecules,
N containing one atom, or form a complex molecule N_{3}, like
oxygen in passing into ozone. Probably under a silent discharge
the molecules of oxygen, O_{2}, are partly decomposed and the
individual atoms O combine with O_{2}, forming ozone, O_{3}.

[13] This reaction, discovered by Chabrié and investigated by Thénard,
was only rightly understood when Deville applied the principles
of dissociation to it.

[14] The action of nitrogen on acetylene (Berthelot) resembles this
reaction. A mixture of these gases under the influence of a
silent discharge gives hydrocyanic acid, C_{2}H_{2} + N_{2} =
2CNH. This reaction cannot proceed beyond a certain limit because
it is reversible.

[15] Berthelot successfully employed electricity of even feeble
potential in these experiments, which fact led him to think that
in nature, where the action of electricity takes place very
frequently, a part of the complex nitrogenous substances may
proceed from the gaseous nitrogen of the air by this method.

As the nitrogenous substances of organisms play a very important
part in them (organic life cannot exist without them), and as
the nitrogenous substances introduced into the soil are capable
of invigorating its crops (of course in the presence of the
other nourishing principles required by plants), the question
of the means of converting the atmospheric nitrogen into the
nitrogenous compounds of the soil, or into _assimilable nitrogen_
capable of being absorbed by plants and of forming complex
(albuminous) substances in them, is one of great theoretical and
practical interest. The artificial (technical) conversion of the
atmospheric nitrogen into nitrogenous compounds, notwithstanding
repeated attempts, cannot yet be considered as fulfilled in a
practical remunerative manner although its possibility is already
evident. Electricity will probably aid in solving this very
important practical problem. When the theoretical side of the
question is further advanced, then without doubt an advantageous
means will be found for the manufacture of nitrogenous substances
from the nitrogen of the air; and this is needed, before all,
for the agriculturist, to whom nitrogenous fertilisers form an
expensive item, and are more important than all other manures.

One thousand tons of farmyard manure do not generally contain
more than four tons of nitrogen in the form of complex
nitrogenous substances, and this amount of nitrogen is contained
in twenty tons of ammonium sulphate, therefore the effect of
a mass of farmyard manure in respect to the introduction of
nitrogen may be produced by small quantities of artificial
nitrogenous fertilisers (_see_ Note 15 bis).

By such indirect methods does the gaseous nitrogen of the atmosphere yield its primary compounds, in which form it enters into plants, and is elaborated in them into complex albuminous substances.[15 bis] But, starting from a given compound of nitrogen with hydrogen or oxygen, we may, without the aid of organisms, obtain, as will afterwards be partially indicated, most diverse and complex nitrogenous substances, which cannot by any means be formed directly from gaseous nitrogen. In this we see an example not only of the difference between an element in the free state and an intrinsic element, but also of those circuitous or _indirect methods_ by which substances are formed in nature. The discovery, prognostication, and, in general, the study of such indirect methods of the preparation and formation of substances forms one of the existing problems of chemistry. From the fact that A does not act at all on B, it must not be concluded that a compound AB is not to be formed. The substances A and B contain atoms which occur in AB, but their state or the nature of their motion may not be at all that which is required for the formation of AB, and in this substance the chemical state of the elements may be as different as the state of the atoms of oxygen in ozone and in water. Thus free nitrogen is inactive; but in its compounds it very easily enters into changes and is distinguished by great activity. An acquaintance with the compounds of nitrogen confirms this. But, before entering on this subject, let us consider air as a mass containing free nitrogen.

[15 bis] Although the numerous, and as far as possible accurate and
varied researches made in the physiology of plants have proved
that the higher forms of plants are not capable of directly
absorbing the nitrogen of the atmosphere and converting it into
complex albuminous substances, still it has been long and
repeatedly observed that the amount of nitrogenous substances in
the soil is increased by the cultivation of plants of the bean
(leguminous) family such as pea, acacia, &c. A closer study of
these plants has shown that this is connected with the formation
of peculiar nodular swellings in their roots caused by the growth
of peculiar micro-organisms (bacteria) which cohabit the soil with
the roots, and are capable of absorbing nitrogen from the air,
_i.e._ of converting it into assimilated nitrogen. This branch of
plant physiology, which forms another proof of the important part
played by micro-organisms in nature, cannot be discussed in this
work, but it should be mentioned, since it is of great theoretical
and practical interest, and, moreover, phenomena of this kind,
which have recently been discovered, promise to explain, to some
extent at least, certain of the complex problems concerning the
development of life on the earth.

Judging from what has been already stated, it will be evident that _atmospheric air_[16] contains a mixture of several gases and vapours. Some of them are met with in it in nearly constant proportions, whilst others, on the contrary, are very variable in their amount. The chief component parts of air, placed in the order of their relative amounts, are the following: nitrogen,[16 bis] oxygen, aqueous vapour, carbonic anhydride, nitric acid, salts of ammonia, oxides of nitrogen, and also ozone, hydrogen peroxide, and complex organic nitrogenous substances. Besides these, air generally contains water, as spray, drops, and snow, and particles of solids, perhaps of cosmic origin in certain instances, but in the majority of cases proceeding from the mechanical translation of solid particles from one locality to another by the wind. These small solid and liquid particles (having a large surface in proportion to their weight) are suspended in air as solid matter is suspended in turbid water; they often settle on the surface of the earth, but the air is never entirely free from them because they are never in a state of complete rest. Then, air not unfrequently contains incidental traces of various substances as everyone knows by experience. These incidental substances sometimes belong to the order of those which act injuriously, the germs of lower organisms--for instance of moulds--and the class of carriers of infectious diseases.

[16] Under the name of atmospheric air the chemist and physicist
understand ordinary air containing nitrogen and oxygen only,
notwithstanding that the other component parts of air have a very
important influence on the living matter of the earth's surface.
That air is so represented in science is based on the fact that
only the two components above-named are met with in air in a
constant quantity, whilst the others are variable. The solid
impurities may be separated from air required for chemical or
physical research by simple filtration through a long layer of
cotton-wool placed in a tube. Organic impurities are removed by
passing the air through a solution of potassium permanganate. The
carbonic anhydride contained in air is absorbed by alkalis--best
of all, soda-lime, which in a dry state in porous lumps absorbs
it with exceeding rapidity and completeness. Aqueous vapour
is removed by passing the air over calcium chloride, strong
sulphuric acid, or phosphoric anhydride. Air thus purified is
accepted as containing only nitrogen and oxygen, although in
reality it still contains a certain quantity of hydrogen and
hydrocarbons, from which it may be purified by passing over
copper oxide heated to redness. The copper oxide then oxidises
the hydrogen and hydrocarbons--it burns them, forming water and
carbonic anhydride, which may be removed as above described. When
it is said that in the determination of the density of gases the
weight of air is taken as unity, it is understood to be such air,
containing only nitrogen and oxygen.

[16 bis] Thanks to the remarkable discovery made in the summer of 1894
by Lord Rayleigh and Prof. Ramsay, the well-known component
elements of air must now he supplemented by 1 p.c. (by volume) of
a heavy gas (density about 19, H = 1), inactive like nitrogen,
which was discovered in the researches made by Lord Rayleigh on
the density of nitrogen as mentioned in note 4 bis. Up to the
present time this gas has been always determined together with
nitrogen, because it combines with neither the hydrogen in the
eudiometer nor with the copper in the gravimetric method of
determining the composition of air, and therefore has always
remained with the nitrogen. It has been possible to separate it
from nitrogen since magnesium absorbs nitrogen at a red heat,
while this gas remains unabsorbed, and was found to have a density
nearly one and a half time greater than that of nitrogen (is it
not a polymer of nitrogen, N_{3}?). It is now known also that this
gas gives a luminous spectrum, which contains the bright blue line
observed in the spectrum of nitrogen. Owing to the fact that it is
an exceedingly inert substance, even more so than nitrogen, it has
been termed Argon. Further reference will be made to it in the
Appendix.

In the air of the various countries of the earth, at different longitudes and at different altitudes above its surface, on the ocean or on the dry land--in a word, in the air of most diverse localities of the earth--the oxygen and nitrogen are found everywhere to be in a constant ratio. This is, moreover, self-evident from the fact that the air constantly diffuses (intermixes by virtue of the internal motion of the gaseous particles) and is also put into motion and intermixed by the wind, by which processes it is equalised in its composition over the entire surface of the earth. In those localities where the air is subject to change, and is in a more or less enclosed space, or, at any rate, in an unventilated space, it may alter very considerably in its composition. For this reason the air in dwellings, cellars, and wells, in which there are substances absorbing oxygen, contains less of this gas, whilst the air on the surface of standing water, which abounds in the lower orders of plant life evolving oxygen, contains an excess of this gas.[17] The constant composition of air over the whole surface of the earth has been proved by a number of most careful researches.[18]

[17] As a further proof of the fact that certain circumstances may
change the composition of air, it will be enough to point out
that the air contained in the cavities of glaciers contains only
up to 10 p.c. of oxygen. This depends on the fact that at low
temperatures oxygen is much more soluble in snow-water and snow
than nitrogen. When shaken up with water the composition of air
should change, because the water dissolves unequal quantities of
oxygen and nitrogen. We have already seen (Chapter I.) that the
air boiled off from water saturated at about 0° contains about
thirty-five volumes of oxygen and sixty-five volumes of nitrogen,
and we have considered the reason of this.

[18] The analysis of air by weight conducted by Dumas and Boussingault
in Paris, which they repeated many times between April 27 and
September 22, 1841, under various conditions of weather, showed
that the amount by weight of oxygen only varies between 22·89
p.c. and 23·08 p.c., the average amount being 23·07 p.c. Brunner,
at Bern in Switzerland, and Bravais, at Faulhorn in the Bernese
Alps, at a height of two kilometres above the level of the sea,
Marignac at Geneva, Lewy at Copenhagen, and Stas at Brussels,
have analysed the air by the same methods, and found that its
composition does not exceed the limits determined for Paris.
The most recent determinations (with an accuracy of ±0·05 p.c.)
confirm the conclusion that the composition of the atmosphere is
constant.

As there are some grounds (which will be mentioned shortly) for
considering that the composition of the air at great altitudes
is slightly different from that at attainable heights--namely,
that it is richer in the lighter nitrogen--several fragmentary
observations made at Munich (Jolly, 1880) gave reason for
thinking that in the upward currents (that is in the region of
minimum barometric pressure or at the centres of meteorological
cyclones) the air is richer in oxygen than in the descending
currents of air (in the regions of anticyclones or of barometric
maxima); but more carefully conducted observations showed
this supposition to be incorrect. Improved methods for the
analysis of air have shown that certain slight variations in
its composition do actually occur, but in the first place they
depend on incidental local influences (on the passage of the air
over mountains and large surfaces of water, regions of forest
and herbage, and the like), and in the second place are limited
to quantities which are scarcely distinguishable from possible
errors in the analyses. The researches made by Kreisler in
Germany (1885) are particularly convincing.

The considerations which lead to the supposition that the
atmosphere at great altitudes contains less oxygen than at the
surface of the earth are based on the law of partial pressures
(Chapter I.) According to this law, the equilibrium of the
oxygen in the strata of the atmosphere is not dependent on
the equilibrium of the nitrogen, and the variation in the
densities of both gases with the height is determined by the
pressure of each gas separately. Details of the calculations
and considerations here involved are contained in my work _On
Barometric Levellings_, 1876, p. 48.

On the basis of the law of partial pressure and of hypsometrical
formulæ, expressing the laws of the variation of pressures at
different altitudes, the conclusion may be deduced that at the
upper strata of the atmosphere the proportion of the nitrogen
with respect to the oxygen increases, but the increase will not
exceed a fraction per cent., even at altitudes of four and a half
to six miles, the greatest height within the reach of men either
by climbing mountains or by means of balloons. This conclusion is
confirmed by the analyses of air collected by Welch in England
during his aëronautic ascents.

_The analysis of air_ is effected by converting the oxygen into a non-gaseous compound, so as to separate it from the air. The original volume of the air is first measured, and then the volume of the remaining nitrogen. The quantity of oxygen is calculated either from the difference between these volumes or by the weight of the oxygen compound formed. All the volumetric measurements have to be corrected for pressure, temperature, and moisture (Chapters I. and II.) The medium employed for converting the oxygen into a non-gaseous substance should enable it to be taken up from the nitrogen to the very end without evolving any gaseous substance. So, for instance,[19] a mixture of pyrogallol, C_{6}H_{6}O_{3}, with a solution of a caustic alkali absorbs oxygen with great ease at the ordinary temperature (the solution turns black), but it is unsuited for accurate analysis because it requires an aqueous solution of an alkali, and it alters the composition of the air by acting on it as a solvent.[20] However, for approximate determinations this simple method gives results which are entirely satisfactory.

[19] The complete absorption of the oxygen may be attained by
introducing moist phosphorus into a definite volume of air; the
occurrence of this is recognised by the fact of the phosphorus
becoming non-luminous in the dark. The amount of oxygen may be
determined by measuring the volume of nitrogen remaining. This
method however cannot give accurate results, owing to a portion
of the air being dissolved in the water, to the combination
of some of the nitrogen with oxygen and to the necessity of
introducing and withdrawing the phosphorus, which cannot be
accomplished without introducing bubbles of air.

[20] For rapid and approximate analyses (technical and hygienic), such
a mixture is very suitable for determining the amount of oxygen
in mixtures of gases from which the substances absorbed by
alkalis have first been removed. According to certain observers,
this mixture evolves a certain (small) quantity of carbonic oxide
after absorbing oxygen.

The determinations in a eudiometer (Chapter III.) give more exact results, if all the necessary corrections for changes of pressure, temperature, and moisture be taken into account. This determination is carried out essentially as follows:--A certain amount of air is introduced into the eudiometer, and its volume is determined. About an equal volume of dry hydrogen is then passed into the eudiometer, and the volume again determined. The mixture is then exploded, in the way described for the determination of the composition of water. The remaining volume of the gaseous mixture is again measured; it will be less than the second of the previously measured volumes. Out of three volumes which have disappeared, one belonged to the oxygen and two to the hydrogen, consequently one-third of the loss of volume indicates the amount of oxygen contained in the air.[21]

[21] Details of eudiometrical analysis must, as was pointed out in
Chap. III., Note 32, be looked for in works on analytical
chemistry. The same remark applies to the other analytical
methods mentioned in this work. They are only described for the
purpose of showing the diversity of the methods of chemical
research.

The most complete method for the analysis of air, and one which is accompanied by the least amount of error, consists in the direct weighing, as far as is possible, of the oxygen, nitrogen, water, and carbonic anhydride contained in it. For this purpose the air is first passed through an apparatus for retaining the moisture and carbonic anhydride (which will be considered presently), and is then led through a tube which contains shavings of metallic copper and has been previously weighed. A long layer of such copper heated to redness absorbs all the oxygen from the air, and leaves pure nitrogen, whose weight must be determined. This is done by collecting it in a weighed and exhausted globe, while the amount by weight of oxygen is shown by the increase in weight of the tube with the copper after the experiment.

Air free from moisture and carbonic anhydride[22] contains 20·95 to 20·88[23] parts by volume of oxygen; the mean amount of oxygen will therefore be 20·92 ± 0·05 per cent. Taking the density of air = 1 and of oxygen = 1·105 and nitrogen 0·972 the composition of air by weight will be 23·12 per cent. of oxygen and 76·88 per cent. of nitrogen.[24]

[22] Air free from carbonic anhydride indicates after explosion the
presence of a small quantity of carbonic anhydride, as De
Saussure remarked, and air free from moisture, after being
passed over red-hot copper oxide, appears invariably to contain
a small quantity of water, as Boussingault has observed. These
observations lead to the assumption that air always contains a
certain quantity of gaseous hydrocarbons, like marsh gas, which,
as we shall afterwards learn, is evolved from the earth, marshes,
&c. Its amount, however, does not exceed a few hundredths per
cent.

[23] The analyses of air are accompanied by errors, and there are
variations of composition attaining hundredths per cent.; the
average normal composition of air is therefore only correct to
the first decimal place.

[24] These figures express the mean composition of air from an average
of the most accurate determinations; they are accurate within
±0·05 p.c.

The possibility of the composition of air being altered by the mere action of a solvent very clearly shows that the component parts of air are in a state of mixture, in which any gases may occur; they do not in this case form a definite compound, although the composition of the atmosphere does appear constant under ordinary conditions. The fact that its composition varies under different conditions confirms the truth of this conclusion, and therefore the constancy of the composition of air must not be considered as in any way dependent on the nature of the gases entering into its composition, but only as proceeding from cosmic phenomena co-operating towards this constancy. It must be admitted, therefore, that the processes evolving oxygen, and chiefly the processes of the respiration of plants, are of equal force with those processes which absorb oxygen over the entire surface of the earth.[25]

[25] In Chapter III., Note 4, an approximate calculation is made for
the determination of the balance of oxygen in the entire
atmosphere; it may therefore he supposed that the composition of
air will vary from time to time, the relation between vegetation
and the oxygen absorbing processes changes; but as the atmosphere
of the earth can hardly have a definite limit and we have
already seen (Chapter IV., Note 33) that there are observations
confirming this, it follows that our atmosphere should vary in
its component parts with the entire heavenly space, and therefore
it must he supposed that any variation in the composition by
weight of the air can only take place exceedingly slowly, and in
a manner imperceptible by experiment.

Air always contains more or less moisture[26] and _carbonic anhydride_ produced by the respiration of animals and the combustion of carbon and carboniferous compounds. The latter shows the properties of an acid anhydride. In order to determine the amount of carbonic anhydride in air, substances are employed which absorb it--namely, alkalis either in solution or solid. A solution of caustic potash, KHO, is poured into light glass vessels, through which the air is passed, and the amount of carbonic anhydride is determined by the increase in weight of the vessel. But it is best to take a solid porous alkaline mass such as soda-lime.[27] With a slow current of air a layer of soda-lime 20 cm. in length is sufficient to completely deprive 1 cubic metre of air of the carbonic anhydride it contains. A series of tubes containing calcium chloride for absorbing the moisture[28] is placed before the apparatus for the absorption of the carbonic anhydride, and a measured mass of air is passed through the whole apparatus by means of an aspirator. In this manner the determination of the moisture is combined with the absorption of the carbonic anhydride. The arrangement shown in fig. 38 is such a combination.

[26] The amount of moisture contained in the air is considered in
greater detail in the study of physics and meteorology and the
subject has been mentioned above, in Chapter I., Note 1, where
the methods of absorbing moisture from gases were pointed out.

[27] Soda-lime is prepared in the following manner:--Unslaked lime
is finely powdered and mixed with a slightly warmed and very
strong solution of caustic soda. The mixing should be done in
an iron dish, and the materials should be well stirred together
until the lime begins to slake. When the mass becomes hot, it
boils, swells up, and solidifies, forming a porous mass very rich
in alkali and capable of rapidly absorbing carbonic anhydride. A
lump of caustic soda or potash presents a much smaller surface
for absorption and therefore acts much less rapidly. It is
necessary to place an apparatus for absorbing water after the
apparatus for absorbing the carbonic anhydride, because the
alkali in absorbing the latter gives off water.

[28] It is evident that the calcium chloride employed for absorbing the
water should be free from lime or other alkalis in order that it
may not retain carbonic anhydride. Such calcium chloride may be
prepared in the following manner: A perfectly neutral solution of
calcium chloride is prepared from lime and hydrochloric acid; it
is then carefully evaporated first on a water-bath and then on a
sand-bath. When the solution attains a certain strength a scum is
formed, which solidifies at the surface. This scum is collected,
and will be found to be free from caustic alkalis. It is
necessary in any case to test it before use, as otherwise a large
error may be introduced into the results, owing to the presence
of free alkali (lime). It is best to pass carbonic anhydride
through the tube containing the calcium chloride for some time
before the experiment, in order to saturate any free alkali that
may remain from the decomposition of a portion of the calcium
chloride by water, CaCl_{2} + 2H_{2}O = CaOH_{2}O + 2HCl.

The amount of carbonic anhydride[29] in free air is incomparably more constant than the amount of moisture. The average amount in 100 volumes of dry air is approximately 0·03 volume--that is, 10,000 volumes of air contain about three volumes of carbonic anhydride, most frequently about 2·95 volumes. As the specific gravity of carbonic anhydride referred to air = 1·52, it follows that 100 parts by weight of air contain 0·045 part by weight of carbonic anhydride. This quantity varies according to the time of year (more in winter), the altitude above the level of the sea (less at high altitudes), the proximity to forests and fields (less) or cities (greater), &c. But the variation is small and rarely exceeds the limits of 2-1/2 to 4 ten-thousandths by volume.[30] As there are many natural local influences which either increase the amount of carbonic anhydride in the air (respiration, combustion, decomposition, volcanic eruptions, &c.), or diminish it (absorption by plants and water), the reason of the great constancy in the amount of this gas in the air must be looked for, in the first place, in the fact that the wind mixes the air of various localities together, and, in the second place, in the fact that the waters of the ocean, holding carbonic acid in solution,[31] form an immense reservoir for regulating the amount of this gas in the atmosphere. Immediately the partial pressure of the carbonic anhydride in the air decreases, the water evolves it, and when the partial pressure increases, it absorbs it, and thus nature supplies the conditions for a natural state of moving equilibrium in this as in so many other instances.[32]

[29] Recourse is had to special methods when the determination only
takes note of the carbonic anhydride of the air. For instance,
it is absorbed by an alkali which does not contain carbonates
(by a solution of baryta or caustic soda mixed with baryta),
and then the carbonic anhydride is expelled by an excess of an
acid, and its amount determined by the volume given off. A rapid
method of determining CO_{2} (for hygienic purposes) is given by
the fall of tension produced by the introduction of an alkali
(the air having been either brought to dryness or saturated
with moisture). Dr. Schidloffsky's apparatus is based upon this
principle. The question as to the amount of carbonic anhydride
present in the air has been submitted to many voluminous and
exact researches, especially those of Reiset, Schloesing, Müntz,
and Aubin, who showed that the amount is not subject to such
variations as at first announced on the basis of incomplete and
insufficiently accurate determinations.

[30] It is a different case in enclosed spaces, in dwellings, cellars,
wells, caves, and mines, where the renewal of air is impeded.
Under these circumstances large quantities of carbonic anhydride
may accumulate. In cities, where there are many conditions for
the evolution of carbonic anhydride (respiration, decomposition,
combustion), its amount is greater than in free air, yet even
in still weather the difference does not often exceed one
ten-thousandth (that is, rarely attains 4 instead of 2·9 vols. in
10000 vols. of air).

[31] In the sea as well as in fresh water, carbonic acid occurs in
two forms, directly dissolved in the water, and combined with
lime as calcium bicarbonate (hard waters sometimes contain very
much carbonic acid in this form). The tension of the carbonic
anhydride in the first form varies with the temperature, and
its amount with the partial pressure, and that in the form of
acid salts is under the same conditions, for direct experiments
have shown a similar dependence in this case, although the
quantitative relations are different in the two cases.

[32] In studying the phenomena of nature the conclusion is arrived at
that the universally reigning state of mobile equilibrium forms
the chief reason for that harmonious order which impresses all
observers. It not unfrequently happens that we do not see the
causes regulating the order and harmony; in the particular
instance of carbonic anhydride, it is a striking circumstance
that in the first instance a search was made for an harmonious
and strict uniformity, and in incidental (insufficiently
accurate and fragmentary) observations conditions were even
found for concluding it to be absent. When, later, the rule of
this uniformity was confirmed, then the causes regulating such
order were also discovered. The researches of Schloesing were of
this character. Deville's idea of the dissociation of the acid
carbonates of sea-water is suggested in them. In many other cases
also, a correct interpretation can only follow from a detailed
investigation.

Besides nitrogen, oxygen, moisture, and carbonic acid, all the other substances occurring in air are found in infinitesimally small quantities by weight, and therefore the _weight of a cubic measure of air_ depends, to a sensible degree, on the above-named components alone. We have already mentioned that at 0° and 760 mm. pressure the weight of a cubic litre of air is 1·293 gram. This weight varies with the acceleration of gravity, _g_, so that if _g_ be expressed in metres the weight of a litre of air, _e_ = _g_ × 0·131844 gram. For St. Petersburg _g_ is about 9·8188, and therefore _e_ is about 1·2946,[33] the air being understood to be dry and free from carbonic anhydride. Taking the amount of the latter as 0·03 per 100 volumes, we obtain a greater weight; for example, for St. Petersburg _e_ = 1·2948 instead of 1·2946 gram. The weight of one litre of moist air in which the tension[34] of the aqueous vapour (partial pressure) = _f_ mm., at a pressure (total) of air of H millimetres, at a temperature _t_, will be (_i.e._, if at 0° and 760 mm. the weight of dry air = _e_) equal to _e_/(1 + 0·00367_t_) × (H - 0·38_f_)/760. For instance, if H = 730 mm., _t_ = 20°, and _f_ = 10 mm. (the moisture is then slightly below 60 p.c.), the weight of a litre of air at St. Petersburg = 1·1527 gram.[35]

[33] The difference of the weight of a litre of dry air (free from
carbonic anhydride) at 0° and 760 mm., at different longitudes
and altitudes, depends on the fact that the force of gravity
varies under these conditions, and with it the pressure of the
barometrical column also varies. This is treated in detail
in my works _On the Elasticity of Gases_ and _On Barometric
Levellings_, and 'The Publications of the Weights and Measures
Department' (_Journal of the Russian Physico-Chemical Society_,
1894).