Chapter III: Oxygen and the Chief Aspects of Its Saline Combinations (1)

On the earth's surface there is no other element which is so widely distributed as oxygen in its various compounds.[1] It makes up eight-ninths of the weight of water, which occupies the greater part of the earth's surface. Nearly all earthy substances and rocks consist of compounds of oxygen with metals and other elements. Thus, the greater part of sand is formed of silica, SiO_{2}, which contains 53 p.c. of oxygen; clay contains water, alumina (formed of aluminium and oxygen), and silica. It may be considered that earthy substances and rocks contain up to one-third of their weight of oxygen; animal and vegetable substances are also very rich in oxygen. Without counting the water present in them, plants contain up to 40, and animals up to 20 p.c. by weight of oxygen. Thus, oxygen compounds predominate on the earth's surface. Besides this, a portion exists in a free state, and is contained in admixture with nitrogen in the atmosphere, forming about one-fourth of its mass, or one-fifth of its volume.

[1] As regards the interior of the earth, it probably contains far less
oxygen compounds than the surface, judging by the accumulated
evidences of the earth's origin, of meteorites, of the earth's
density, &c. (_see_ Chapter VIII., Note 58, and Chapter XXII.,
Note 2).

Being so widely distributed in nature, oxygen plays a very important part in it, for a number of the phenomena which take place before us are mainly dependent on it. _Animals breathe_ air in order to obtain only _oxygen_ from it, the oxygen entering into their respiratory organs (the lungs of human beings and animals, the gills of fishes, and the trachæ of insects); they, so to say, drink in air in order to absorb the oxygen. The oxygen of the air (or dissolved in water) passes through the membranes of the respiratory organs into the blood, is retained in it by the blood corpuscles, is transmitted by their means to all parts of the body, aids their transformations, bringing about chemical processes in them, and chiefly extracting carbon from them in the form of carbonic anhydride, the greater part of which passes into the blood, is dissolved by it, and is thrown off by the lungs during the absorption of the oxygen. Thus, in the process of respiration carbonic anhydride (and water) is given off, and the oxygen of the air absorbed, by which means the blood is changed from a red venous to a dark-red arterial blood. The cessation of this process causes death, because then all those chemical processes, and the consequent heat and work which the oxygen introduced into the system brought about, cease. For this reason suffocation and death ensue in a vacuum, or in a gas which does not contain free oxygen, _i.e._ which does not support combustion. If an animal be placed in an atmosphere of free oxygen, at first its movements are very active and a general invigoration is remarked, but a reaction soon sets in, and death may ensue. The oxygen of the air when it enters the lungs is diluted with four volumes of nitrogen, which is not absorbed into the system, so that the blood absorbs but a small quantity of oxygen from the air, whilst in an atmosphere of pure oxygen a large quantity of oxygen would be absorbed, and would produce a very rapid change of all parts of the organism, and destroy it. From what has been said, it will be understood that oxygen may be employed in respiration, at any rate for a limited time, when the respiratory organs suffer under certain forms of suffocation and impediment to breathing.[2]

[2] It is evident that the partial pressure (_see_ Chapter I.) acts in
respiration. The researches of Paul Bert showed this with
particular clearness. Under a pressure of one-fifth of an
atmosphere consisting of oxygen only, animals and human beings
remain under the ordinary conditions of the partial pressure
of oxygen, but organisms cannot support air rarefied to
one-fifth, for then the partial pressure of the oxygen falls
to one-twenty-fifth of an atmosphere. Even under a pressure of
one-third of an atmosphere the regular life of human beings
is impossible, by reason of the impossibility of respiration
(because of the decrease of solubility of oxygen in the blood),
owing to the small partial pressure of the oxygen, and not from
any mechanical effect of the decrease of pressure. Paul Bert
illustrated all this by many experiments, some of which he
conducted on himself. This explains, among other things, the
discomfort felt in the ascent of high mountains or in balloons
when the height reached exceeds eight kilometres, and at pressures
below 250 mm. (Chapter II., Note 23). It is evident that an
artificial atmosphere has to be employed in the ascent to great
heights, just as in submarine work. The cure by compressed and
rarefied air which is practised in certain illnesses is based
partly on the mechanical action of the change of pressure, and
partly on the alteration in the partial pressure of the respired
oxygen.

The combustion of organic substances--that is, substances which make up the composition of plants and animals--proceeds in the same manner as the combustion of many inorganic substances, such as sulphur, phosphorus, iron, &c., from the combination of these substances with oxygen, as was described in the Introduction. The decomposition, rotting, and similar transformations of substances, which proceed around us, are also very often dependent on the action of the oxygen of the air, and also reduce it from a free to a combined state. The majority of the compounds of oxygen are, like water, very stable, and do not give up their oxygen under the ordinary conditions of nature. As these processes are taking place everywhere, it might be expected that the amount of free oxygen in the atmosphere should decrease, and this decrease should proceed somewhat rapidly. This is, in fact, observed where combustion or respiration proceeds in a closed space. Animals suffocate in a closed space because in consuming the oxygen the air remains unfit for respiration. In the same manner combustion, after a time, ceases in a closed space, which may be proved by a very simple experiment. An ignited substance--for instance, a piece of burning sulphur--has only to be placed in a glass flask, which is then closed with a stout cork to prevent the access of the external air; combustion will proceed for a certain time, so long as the flask contains any free oxygen, but it will cease when the oxygen of the enclosed air has combined with the sulphur. From what has been said, it is evident that regularity of combustion or respiration requires a constant renewal of air--that is, that the burning substance or respiring animal should have access to a fresh supply of oxygen. This is attained in dwellings by having many windows, outlets, and ventilators, and by the current of air produced by fires and stoves. As regards the air over the entire earth's surface its amount of oxygen hardly decreases, because in nature there is a process going on which renews the supply of free oxygen. _Plants_, or rather their leaves, during daytime,[3] under the influence of light, absorb carbonic anhydride CO_{2}, and _evolve free oxygen_. Thus the loss of oxygen which occurs in consequence of the respiration of animals and of combustion is made good by plants. If a leaf be placed in a bell jar containing water, and carbonic anhydride (because this gas is absorbed and oxygen evolved from it by plants) be passed into the bell, and the whole apparatus placed in sunlight, then oxygen will accumulate in the bell jar. This experiment was first made by Priestley at the end of the last century. Thus the life of plants on the earth not only serves for the formation of food for animals, but also for keeping up a constant percentage of oxygen in the atmosphere. In the long period of the life of the earth an equilibrium has been attained between the processes absorbing and evolving oxygen, by which a definite quantity of free oxygen is preserved in the entire mass of the atmosphere.[4]

[3] At night, without the action of light, without the absorption
of that energy which is required for the decomposition of carbonic
anhydride into free oxygen and carbon (which is retained by the
plants) they breathe like animals, absorbing oxygen and evolving
carbonic anhydride. This process also goes on side by side with
the reverse process in the daytime, but it is then far feebler
than that which gives oxygen.

[4] The earth's surface is equal to about 510 million square
kilometres, and the mass of the air (at a pressure of 760 mm.)
on each kilometre of surface is about 10-1/3 thousand millions
of kilograms, or about 10-1/3 million tons; therefore the whole
weight of the atmosphere is about 5,100 million million (= 51 ×
10^{14}) tons. Consequently there are about 2 × 10^{15} tons of
free oxygen in the earth's atmosphere. The innumerable series of
processes which absorb a portion of this oxygen are compensated
for by the plant processes. Assuming that 100 million tons of
vegetable matter, containing 40 p.c. of carbon, formed from
carbonic acid, are produced (and the same process proceeds in
water) per year on the 100 million square kilometres of dry land
(ten tons of roots, leaves, stems, &c., per hectare, or 1/100 of
a square kilometre), we find that the plant life of the dry land
gives about 100,000 tons of oxygen, which is an insignificant
fraction of the entire mass of the oxygen of the air.

Oxygen was obtained as an independent gas in 1774 by Priestley in England and in the same year by Scheele in Sweden, but its nature and great importance were only perfectly elucidated by Lavoisier.

Free oxygen may be obtained by one or other method from all the substances in which it occurs. Thus, for instance, the oxygen of many substances may be transferred into water, from which, as we have already seen, oxygen may be obtained.[5] We will first consider the methods of extracting oxygen from air as being a substance everywhere distributed. The separation of oxygen from it is, however, hampered by many difficulties.

[5] The extraction of oxygen from water may be effected by two
processes: either by the decomposition of water into its
constituent parts by the action of a galvanic current (Chapter
II.), or by means of the removal of the hydrogen from water.
But, as we have seen and already know, hydrogen enters into
direct combination with very few substances, and then only under
special circumstances; whilst oxygen, as we shall soon learn,
combines with nearly all substances. Only gaseous chlorine (and,
especially, fluorine) is capable of decomposing water, taking up
the hydrogen from it, without combining with the oxygen. Chlorine
is soluble in water, and if an aqueous solution of chlorine,
so-called chlorine water, be poured into a flask, and this flask
be inverted in a basin containing the same chlorine water, then we
shall have an apparatus by means of which oxygen may be extracted
from water. At the ordinary temperature, and in the dark, chlorine
does not act on water, or only acts very feebly; but under the
action of direct sunlight chlorine decomposes water, with the
evolution of oxygen. The chlorine then combines with the hydrogen,
and gives hydrochloric acid, which dissolves in the water, and
therefore free oxygen only will be separated from the liquid, and
it will only contain a small quantity of chlorine in admixture,
which can be easily removed by passing the gas through a solution
of caustic potash.

From air, which contains a _mixture_ of oxygen and nitrogen, the nitrogen alone cannot be removed, because it has no inclination to combine directly or readily with any substance; and although it does combine with certain substances (boron, titanium), these substances combine simultaneously with the oxygen of the atmosphere.[6] However, oxygen may be separated from air by causing it to combine with substances which may be easily decomposed by the action of heat, and, in so doing, give up the oxygen absorbed--that is, by making use of reversible reactions. Thus, for instance, the oxygen of the atmosphere may be made to oxidise sulphurous anhydride, SO_{2} (by passing directly over ignited spongy platinum), and to form sulphuric anhydride, or sulphur trioxide, SO_{3}; and this substance (which is a solid and volatile, and therefore easily separated from the nitrogen and sulphurous anhydride), on further heating, gives oxygen and sulphurous anhydride. Caustic soda or lime extracts (absorbs) the sulphurous anhydride from this mixture, whilst the oxygen is not absorbed, and thus it is isolated from the air. On a large scale in works, as we shall afterwards see, sulphurous anhydride is transformed into hydrate of sulphuric trioxide, or sulphuric acid, H_{2}SO_{4}; if this is allowed to drop on to red-hot flagstones, water, sulphurous anhydride, and oxygen are obtained. The oxygen is easily isolated from this mixture by passing the gases over lime. The extraction of oxygen from oxide of mercury (Priestley, Lavoisier), which is obtained from mercury and the oxygen of the atmosphere, is also a reversible reaction by which oxygen may be obtained from the atmosphere. So also, by passing dry air through a red-hot tube containing barium oxide, it is made to combine with the oxygen of the air. In this reaction the so-called barium peroxide, BaO_{2}, is formed from the barium oxide, BaO, and at a higher temperature the former evolves the absorbed oxygen, and leaves the barium oxide originally taken.[7]

[6] A difference in the physical properties of both gases cannot be
here taken advantage of, because they are very similar in this
respect. Thus the density of oxygen is 16 times and of nitrogen
14 times greater than the density of hydrogen, and therefore
porous vessels cannot be here employed--the difference between
the times of their passage through a porous surface would be too
insignificant.

Graham, however, succeeded in enriching air in oxygen by passing
it through india-rubber. This may be done in the following
way:--A common india-rubber cushion, E (Fig. 27), is taken, and
its orifice hermetically connected with an air-pump, or, better
still, a mercury aspirator (the Sprengel pump is designated by
the letters A, C, B). When the aspirator (Chapter II., Note 16)
has pumped out the air, which will be seen by the mercury running
out in an almost uninterrupted stream, and from its standing
approximately at the barometric height, then it may be clearly
observed that gas passes through the india-rubber. This is also
seen from the fact that bubbles of gas continually pass along
with the mercury. A minus pressure may be constantly maintained
in the cushion by pouring mercury into the funnel A, and screwing
up the pinchcock C, so that the stream flowing from it is small,
and then a portion of the air passing through the india-rubber
will be carried along with the mercury. This air may be collected
in the cylinder, R. Its composition proves to be about 42 volumes
of oxygen with 57 volumes of nitrogen, and one volume of carbonic
anhydride, whilst ordinary air contains only 21 volumes of oxygen
in 100 volumes. A square metre of india-rubber surface (of the
usual thickness) passes about 45 c.c. of such air per hour. This
experiment clearly shows that india-rubber is permeable to gases.
This may, by the way, be observed in common toy balloons filled
with coal-gas. They fall after a day or two, not because there
are holes in them, but because air penetrates into, and the gas
from, their interior, through the surface of the india-rubber of
which they are made. The rate of the passage of gases through
india-rubber does not, as Mitchell and Graham showed, depend
on their densities, and consequently its permeability is not
determined by orifices. It more resembles dialysis--that is, the
penetration of liquids through colloid surfaces. Equal volumes
of gases penetrate through india-rubber in periods of time which
are related to each other as follows:--carbonic anhydride, 100;
hydrogen, 247; oxygen, 532; marsh gas, 633; carbonic oxide,
1,220; nitrogen, 1,358. Hence nitrogen penetrates more slowly
than oxygen, and carbonic anhydride more quickly than other
gases. 2·556 volumes of oxygen and 13·585 volumes of carbonic
anhydride penetrate in the same time as one volume of nitrogen.
By multiplying these ratios by the amounts of these gases in air,
we obtain figures which are in almost the same proportion as the
volumes of the gases penetrating from air through india-rubber. If
the process of dialysis be repeated on the air which has already
passed through india-rubber, then a mixture containing 65 p.c. by
volume of oxygen is obtained. It may be thought that the cause
of this phenomenon is the absorption or occlusion (_see_ Chap.
II., Note 37) of gases by india-rubber and the evolution of the
gas dissolved in a vacuum; and, indeed, india-rubber does absorb
gases, especially carbonic anhydride. Graham called the above
method of the decomposition of air _atmolysis_.

[7] The preparation of oxygen by this method, which is due to
Boussingault, is conducted in a porcelain tube, which is placed
in a stove heated by charcoal, so that its ends project beyond
the stove. Barium oxide (which may be obtained by igniting barium
nitrate, previously dried) is placed in the tube, one end of
which is connected with a pair of bellows, or a gas-holder, for
keeping up a current of air through it. The air is previously
passed through a solution of caustic potash, to remove all
traces of carbonic anhydride, and it is very carefully dried
(for the hydrate BaH_{2}O_{2} does not give the peroxide). At a
_dark-red heat_ (500-600°) the oxide of barium absorbs oxygen
from the air, so that the gas leaving the tube consists almost
entirely of nitrogen. When the absorption ceases, the air will
pass through the tube unchanged, which may be recognised from the
fact that it supports combustion. The barium oxide is converted
into peroxide under these circumstances, and eleven parts of
barium oxide absorb about one part of oxygen by weight. When the
absorption ceases, one end of the tube is closed, a cork with a
gas-conducting tube is fixed into the other end, and the heat of
the stove is increased to a _bright-red heat_ (800°). At this
temperature the barium peroxide gives up all that oxygen which it
acquired at a dark-red heat--_i.e._ about one part by weight of
oxygen is evolved from twelve parts of barium peroxide. After the
evolution of the oxygen there remains the barium oxide which was
originally taken, so that air may be again passed over it, and
thus the preparation of oxygen from one and the same quantity of
barium oxide may be repeated many times. Oxygen has been produced
one hundred times from one mass of oxide by this method; all the
necessary precautions being taken, as regards the temperature of
the mass and the removal of moisture and carbonic acid from the
air. Unless these precautions be taken, the mass of oxide soon
spoils.

As oxygen may become of considerable technical use, from its
capacity for giving high temperatures and intense light in
the combustion of substances, its preparation directly from
air by practical methods forms a problem whose solution many
investigators continue to work at up to the present day. The
most practical methods are those of Tessié du Motay and Kassner.
The first is based on the fact that a mixture of equal weights
of manganese peroxide and caustic soda at an incipient red
heat (about 350°) absorbs oxygen from air, with the separation
of water, according to the equation MnO_{2} + 2NaHO + O =
Na_{2}MnO_{4} + H_{2}O. If superheated steam, at a temperature
of about 450°, be then passed through the mixture, the manganese
peroxide and caustic soda originally taken are regenerated, and
the oxygen held by them is evolved, according to the reverse
equation Na_{2}MnO_{4} + H_{2}O = MnO_{2} + 2NaHO + O. This mode
of preparing oxygen may be repeated for an infinite number of
times. The oxygen in combining liberates water, and steam, acting
on the resultant substance, evolves oxygen. Hence all that is
required for the preparation of oxygen by this method is fuel
and the alternate cutting off the supply of air and steam. In
Kassner's process (1891) a mixture of oxide of lead and lime (PbO
+ 2CaO) is heated to redness in the presence of air, oxygen is
then absorbed and calcium plumbate, Ca_{2}PbO_{4}, formed. The
latter is of a chocolate colour, and on further heating evolves
oxygen and gives the original mixture PbO + 2CaO--that is, the
phenomenon is essentially the same as in Boussingault's process
(with BaO), but according to Le Chatelier (1893) the dissociation
tension of the oxygen evolved from Ca_{2}PbO_{4} is less than with
BaO_{2} at equal temperatures; for instance, at 940°, 112 mm. of
mercury for the first, and for the latter 210 mm. at 720°, and 670
mm. at 790°, while for Ca_{2}PbO_{4} this tension is only reached
at 1,080°. However, in Kassner's process the oxygen is absorbed
more rapidly, and the influence of the presence of moisture and
CO_{2} in the air is not so marked, so that this process, like
that of Tessié du Motay, deserves consideration.

Oxygen is evolved with particular ease by a whole series of unstable oxygen compounds, of which we shall proceed to take a general survey, remarking that many of these reactions, although not all, belong to the number of reversible reactions;[8] so that in order to obtain many of these substances (for instance, potassium chlorate) rich in oxygen, recourse must be had to indirect methods (see Introduction) with which we shall become acquainted in the course of this book.

[8] Even the decomposition of manganese peroxide is reversible, and it
may be re-obtained from that suboxide (or its salts), which
is formed in the evolution of oxygen (Chap. XI., Note 6). The
compounds of chromic acid containing the trioxide CrO_{3} in
evolving oxygen give chromium oxide, Cr_{2}O_{3}, but they re-form
the salt of chromic acid when heated to redness in air with an
alkali.

1. _The compounds of oxygen_ with certain metals, and especially with the so-called noble metals--that is, mercury, silver, gold, and platinum--having once been obtained, retain their oxygen at the ordinary temperature, but part with it at a red heat. The compounds are solids, generally amorphous and infusible, and are easily decomposed by heat into the metal and oxygen. We have seen an example of this in speaking of the decomposition of mercury oxide. Priestley, in 1774, obtained pure oxygen for the first time by heating mercury oxide by means of a burning-glass, and clearly showed its difference from air. He showed its characteristic property of supporting combustion 'with remarkable vigour,' and named it dephlogisticated air.

2. The substances called _peroxides_[9] evolve oxygen at a greater or less heat (and also by the action of many acids). They usually contain metals combined with a large quantity of oxygen. Peroxides are the highest oxides of certain metals; those metals which form them generally give several compounds with oxygen. Those of the lowest degrees of oxidation, containing the least amount of oxygen, are generally substances which are capable of easily reacting with acids--for instance, with sulphuric acid. Such low oxides are called bases. Peroxides contain more oxygen than the bases formed by the same metals. For example, lead oxide contains 7·1 parts of oxygen in 100 parts, and is basic, but lead peroxide contains 13·3 parts of oxygen in 100 parts. _Manganese peroxide_ is a similar substance, and is a solid of a dark colour, which occurs in nature. It is employed for technical purposes under the name of black oxide of manganese (in German, 'Braunstein,' the pyrolusite of the mineralogist). Peroxides are able to evolve oxygen at a more or less elevated temperature. They do not then part with all their oxygen, but with only a portion of it, and are converted into a lower oxide or base. Thus, for example, lead peroxide, on heating, gives oxygen and lead oxide. The decomposition of this peroxide proceeds tolerably easily on heating, even in a glass vessel, but manganese peroxide only evolves oxygen at a strong red heat, and therefore oxygen can only be obtained from it in iron, or other metallic, or clay vessels. This was formerly the method for obtaining oxygen. Manganese peroxide only parts with one-third of its oxygen (according to the equation 3MnO_{2} = Mn_{3}O_{4} + O_{2}), whilst two-thirds remain in the solid substance which forms the residue after heating. Metallic peroxides are also capable of evolving oxygen on heating with sulphuric acid. They then evolve just that amount of oxygen which is in excess of that necessary for the formation of the base, the latter reacting on the sulphuric acid forming a compound (salt) with it. Thus barium peroxide, when heated with sulphuric acid, forms oxygen and barium oxide, which gives a compound with sulphuric acid termed barium sulphate (BaO_{2} + H_{2}SO_{4} = BaSO_{4} + H_{2}O + O).[9 bis] This reaction usually proceeds with greater ease than the decomposition of peroxides by heat alone. For the purposes of experiment powdered manganese peroxide is usually taken and mixed with strong sulphuric acid in a flask, and the apparatus set up as shown in Fig. 28. The gas which is evolved is passed through a Woulfe's bottle containing a solution of caustic potash, to purify it from carbonic anhydride and chlorine, which accompany the evolution of oxygen from commercial manganese peroxide, and the gas is not collected until a thin smouldering taper placed in front of the escape orifice bursts into flame, which shows that the gas coming off is oxygen. By this method of decomposition of the manganese peroxide by sulphuric acid there is evolved, not, as in heating, one-third, but one-half of the oxygen contained in the peroxide (MnO_{2} + H_{2}SO_{4} = MnSO_{4} + H_{2}O + O)--that is, from 50 grams of peroxide about 7-1/5 grams, or about 5-1/2 litres, of oxygen,[10] whilst by heating only about 3-1/2 litres are obtained. The chemists of Lavoisier's time generally obtained oxygen by heating manganese peroxide. At the present time more convenient methods are known.

[9] We shall afterwards see that it is only substances like barium
peroxide (which give hydrogen peroxide) which should be counted
as true peroxides, and that MnO_{2}, PbO_{2}, &c., should be
distinguished from them (they do not give hydrogen peroxide with
acids), and therefore it is best to call them dioxides.

[9 bis] Peroxide of barium also gives oxygen at the ordinary
temperature in the presence of the solutions of many substances
in a higher degree of oxidation. In this respect we may mention
that Kassner (1890) proposes to obtain oxygen for laboratory
purposes by mixing BaO_{2} with FeK_{3}(CN)_{6} (red prussiate of
potash, Chapter XXII.): the reaction proceeds with the evolution
of oxygen even on the addition of a very small quantity of
water. In order to ensure a gradual evolution of gas the author
proposes to introduce both substances into the reaction, little
by little, instead of all at once, which may be done with the
following arrangement (Gavaloffsky): finely powdered peroxide
of barium is placed in an ordinary flask and sufficient water
is added to fill the flask one-third full. The cork closing the
flask has three holes; (1) for the gas-conducting tube; (2) for a
rod to stir the BaO_{2}; and (3) for a glass rod terminating in a
perforated glass vessel containing crystals of FeK_{3}(CN)_{6}.
When it is desired to start the evolution of the oxygen, the
vessel is lowered until it is immersed in the liquid in the flask,
and the BaO_{2} is stirred with the other rod. The reaction
proceeds according to the equation, BaO_{2} + 2FeK_{3}(CN)_{6}
= FeK_{4}(CN)_{6} + FeK_{2}Ba(CN)_{6} + O_{2}. The double salt,
FeBa_{2}(CN)_{6}, crystallises out from the mother liquor. To
understand the course of the reaction, it must be remembered
BaO_{2} is of a higher degree of oxidation, and that it parts
with oxygen and gives the base BaO which enters into the complex
salt FeK_{2}Ba(CN)_{6} = Fe(CN)_{2} + 2KCN + Ba(CN)_{2}, and this
latter = BaO + 2HCN-H_{2}O. Moreover, FeK_{3}(CN)_{6} contains the
salt Fe_{2}(CN)_{6} which also corresponds to the higher degree
of oxidation of iron, Fe_{2}O_{3}, whilst after the reaction a
salt is obtained which contains Fe(CN)_{2}, and corresponds to the
lower degree of oxidation, FeO, so that (in the presence of water)
oxygen is also set free on this side also, _i.e._ the reaction
gives lower degrees of oxidation and oxygen.

[10] Scheele, in 1785, discovered the method of obtaining oxygen by
treating manganese peroxide with sulphuric acid.

3. A third source to which recourse may be had for obtaining oxygen is represented in _acids_ and _salts_ containing much oxygen, which are capable, by parting with a portion or all of their oxygen, of being converted into other compounds (lower products of oxidation) which are more difficultly decomposed. These acids and salts (like peroxides) evolve oxygen either on heating alone, or only when in the presence of some other substance. Sulphuric acid may be taken as an example of an acid which is decomposed by the action of heat alone,[11] for it breaks up at a red heat into water, sulphurous anhydride, and oxygen, as was mentioned before. Priestley, in 1772, and Scheele, somewhat later, obtained oxygen by heating nitre to a red heat. The best examples of the formation of oxygen by the heating of salts is given in _potassium chlorate_, or Berthollet's salt, so called after the French chemist who discovered it. Potassium chlorate is a salt composed of the elements potassium, chlorine, and oxygen, KClO_{3}. It occurs as transparent colourless plates, is soluble in water, especially in hot water, and resembles common table salt in some of its reactions and physical properties; it melts on heating, and in melting begins to decompose, evolving oxygen gas. This decomposition ends in all the oxygen being evolved from the potassium chlorate, potassium chloride being left as a residue, according to the equation KClO_{3} = KCl + O_{3}.[12] This decomposition proceeds at a temperature which allows of its being conducted in a glass vessel. However, in decomposing, the molten potassium chlorate swells up and boils, and gradually solidifies, so the evolution of the oxygen is not regular, and the glass vessel may crack. In order to overcome this inconvenience, the potassium chlorate is crushed and mixed with a powder of a substance which is infusible, incapable of combining with the oxygen evolved, and is a good conductor of heat. Usually it is mixed with manganese peroxide.[13] The decomposition of the potassium chlorate is then considerably facilitated, and proceeds at a lower temperature (because the entire mass is then better heated, both externally and internally), without swelling up, and this method is therefore more convenient than the decomposition of the salt alone. This method for the preparation of oxygen is very convenient; it is generally employed when a small quantity of oxygen is required. Further, potassium chlorate is easily obtained pure, and it evolves much oxygen. 100 grams of the salt give as much as 39 grams, or 30 litres, of oxygen. This method is so simple and easy,[14] that a course of practical chemistry is often commenced by the preparation of oxygen by this method, and of hydrogen by the aid of zinc and sulphuric acid, since by means of these gases many interesting and striking experiments may be performed.[15]

[11] All acids rich in oxygen, and especially those whose elements form
lower oxides, evolve oxygen either directly at the ordinary
temperature (for instance, ferric acid), or on heating (nitric,
manganic, chromic, chloric, and others), or if basic lower oxides
are formed from them, by heating with sulphuric acid. Thus
the salts of chromic acid (for example, potassium dichromate,
K_{2}Cr_{2}O_{7}) give oxygen with sulphuric acid; first
potassium sulphate, K_{2}SO_{4}, is formed, and then the chromic
acid set free gives a sulphuric acid salt of the lower oxide,
Cr_{2}O_{3}.

[12] This reaction is not reversible, and is exothermal--that is, it
does not absorb heat, but, on the contrary, evolves 9,713
calories per molecular weight KClO_{3}, equal to 122 parts of
salt (according to the determination of Thomsen, who burnt
hydrogen in a calorimeter either alone or with a definite
quantity of potassium chlorate mixed with oxide of iron). It does
not proceed at once, but first forms perchlorate, KClO_{4} (_see_
Chlorine and Potassium). It is to be remarked that potassium
chloride melts at 766°, potassium chlorate at 359°, and potassium
perchlorate at 610°. (Concerning the decomposition of KClO_{3},
_see_ Chapter II., Note 47.)

[13] The peroxide does not evolve oxygen in this case. It may be
replaced by many oxides--for instance, by oxide of iron. It is
necessary to take the precaution that no combustible substances
(such as bits of paper, splinters, sulphur, &c.) fall into the
mixture, as they might cause an explosion.

[14] The decomposition of a mixture of fused and well-crushed potassium
chlorate with powdered manganese peroxide proceeds at so low a
temperature (the salt does not melt) that it may be effected in
an ordinary glass flask. The apparatus is arranged in the same
manner as in the decomposition of mercury oxide (Introduction),
or as shown in the last drawing. As the reaction is exothermal,
the decomposition of potassium chlorate with the formation of
oxygen may probably be accomplished, under certain conditions
(for example, under contact action), at very low temperatures.
Substances mixed with the potassium chlorate probably act
partially in this manner.

[15] Many other salts evolve oxygen by heat, like potassium chlorate,
but they only part with it either at a very high temperature (for
instance, common nitre) or else are unsuited for use on account
of their cost (potassium manganate), or evolve impure oxygen at a
high temperature (zinc sulphate at a red heat gives a mixture of
sulphurous anhydride and oxygen), and are not therefore used in
practice.

A solution of _bleaching powder_, which contains calcium hypochlorite, CaCl_{2}O_{2}, evolves oxygen on gently heating when a small quantity of certain oxides is added--for instance, cobalt oxide, which in this case acts by contact (_see_ Introduction). When heated by itself, a solution of bleaching powder does not evolve oxygen, but it oxidises the cobalt oxide to a higher degree of oxidation; this higher oxide of cobalt in contact with the bleaching powder decomposes into oxygen and lower oxidation products, and the resultant lower oxide of cobalt with bleaching powder again gives the higher oxide, which again gives up its oxygen, and so on.[16] The calcium hypochlorite is here decomposed according to the equation CaCl_{2}O_{2} = CaCl_{2} + O_{2}. In this manner a small quantity of cobalt oxide[17] is sufficient for the decomposition of an indefinitely large quantity of bleaching powder.

[16] Such is, at present, the only possible method of explaining the
phenomenon of contact action. In many cases, such as the present
one, it is supported by observations based on facts. Thus,
for instance, it is known, as regards oxygen, that often two
substances rich in oxygen retain it so long as they are separate,
but directly they come into contact free oxygen is evolved from
both of them. Thus, an aqueous solution of hydrogen peroxide
(containing twice as much oxygen as water) acts in this manner on
silver oxide (containing silver and oxygen). This reaction takes
place at the ordinary temperature, and the oxygen is evolved
from both compounds. To this class of phenomena may be also
referred the fact that a mixture of barium peroxide and potassium
manganate with water and sulphuric acid evolves oxygen at the
ordinary temperature (Note 9 bis). It would seem that the essence
of phenomena of this kind is entirely and purely a property of
contact; the distribution of the atoms is changed by contact,
and if the equilibrium be unstable it is destroyed. This is more
especially evident in the case of those substances which change
exothermally--that is, for those reactions which are accompanied
by an evolution of heat. The decomposition CaCl_{2}O_{2} =
CaCl_{2} + O_{2} belongs to this class (like the decomposition of
potassium chlorate).

[17] Generally a solution of bleaching powder is alkaline (contains
free lime), and therefore, a solution of cobalt chloride is added
directly to it, by which means the oxide of cobalt required for
the reaction is formed.

_The properties of oxygen._[18]--It is a permanent _gas_--that is, it cannot be liquefied by pressure at the ordinary temperature, and further, is only liquefied with difficulty (although more easily than hydrogen) at temperatures below -120°, because this is its absolute boiling point. As its critical pressure[19] is about 50 atmospheres, it can be easily liquefied under pressures greater than 50 atmospheres at temperatures below -120°. According to Dewar, the density of oxygen in a critical state is 0·65 (water = 1), but, like all other substances in this state,[20] it varies considerably in density with a change of pressure and temperature, and therefore many investigators who made their observations under high pressures give a greater density, as much as 1·1. Liquefied oxygen is an exceedingly mobile transparent liquid, with a faint blue tint and boiling (tension = 1 atmosphere) about -180°. Oxygen, like all gases, is transparent, and like the majority of gases, colourless. It has no smell or taste, which is evident from the fact of its being a component of air. The weight of one litre of oxygen gas at 0° and 760 mm. pressure is 1·4298 gram; it is therefore slightly denser than air. Its density in respect to air = 1·1056 and in respect to hydrogen = 16.[21]

[18] It must be remarked that in all the reactions above mentioned the
formation of oxygen may be prevented by the admixture of
substances capable of combining with it--for example, charcoal,
many carbon (organic) compounds, sulphur, phosphorus, and various
lower oxidation products, &c. These substances absorb the oxygen
evolved, combine with it, and a compound containing oxygen is
formed. Thus, if a mixture of potassium chlorate and charcoal be
heated, no oxygen is obtained, but an explosion takes place from
the rapid formation of gases resulting from the combination of
the oxygen of the potassium chlorate with the charcoal and the
evolution of gaseous CO_{2}.

The oxygen obtained by any of the above-described methods is
rarely pure. It generally contains aqueous vapour, carbonic
anhydride, and very often small traces of chlorine. The oxygen
may be freed from these impurities by passing it through a
solution of caustic potash, and by drying it. If the potassium
chlorate be dry and pure, it gives almost pure oxygen. However,
if the oxygen be required for respiration in cases of sickness,
it should be washed by passing it through a solution of caustic
alkali and through water. The best way to obtain pure oxygen
directly is to take potassium perchlorate (KClO_{4}), which can
be well purified and then evolves pure oxygen on heating.

[19] With regard to the absolute boiling point, critical pressure, and
the critical state in general, _see_ Chapter II., Notes 29 and 34.

[20] Judging from what has been said in Note 34 of the last chapter,
and also from the results of direct observation, it is evident
that all substances in a critical state have a large coefficient
of expansion, and are very compressible.

[21] As water consists of 1 volume of oxygen and 2 volumes of hydrogen,
and contains 16 parts by weight of oxygen per 2 parts by weight
of hydrogen, it therefore follows directly that oxygen is 16
times denser than hydrogen. Conversely, the composition of water
by weight may be deduced from the densities of hydrogen and
oxygen, and the volumetric composition of water. This method of
mutual and reciprocal correction strengthens the practical data
of the exact sciences, whose conclusions require the greatest
possible exactitude and variety of corrections.

It must he observed that the specific heat of oxygen at constant
pressure is 0·2175, consequently it is to the specific heat of
hydrogen (3·409) as 1 is to 15·6. Hence, the specific heats are
inversely proportional to the weights of equal volumes. This
signifies that equal volumes of both gases have (nearly) equal
specific heats--that is, they require an equal quantity of heat
for raising their temperature by 1°. We shall afterwards consider
the specific heat of different substances more fully in Chap. XIV.

Oxygen, like the majority of difficultly-liquefiable gases, is
but slightly soluble in water and other liquids. The solubility
is given in Note 30, Chap. I. From this it is evident that water
standing in air must absorb--_i.e._ dissolve--oxygen. This oxygen
serves for the respiration of fishes. Fishes cannot exist in
boiled water, because it does not contain the oxygen necessary
for their respiration (_see_ Chap. I.)

In its chemical properties oxygen is remarkable from the fact that it very easily--and, in a chemical sense, vigorously--reacts on a number of substances, forming oxygen compounds. However, only a few substances and mixtures of substances (for example, phosphorus, copper with ammonia, decomposing organic matter, aldehyde, pyrogallol with an alkali, &c.) combine directly with oxygen at the ordinary temperature, whilst many substances easily combine with oxygen at a red heat, and often this combination presents a rapid chemical reaction accompanied by the evolution of a large quantity of heat. Every reaction which takes place rapidly, if it be accompanied by so great an evolution of heat as to produce incandescence, is termed _combustion_. Thus combustion ensues when many metals are plunged into chlorine, or oxide of sodium or barium into carbonic anhydride, or when a spark falls on gunpowder. A great many substances are combustible in oxygen, and, owing to its presence, in air also. In order to start combustion it is generally necessary[22] that the combustible substance should be brought to a state of incandescence. The continuation of the process does not require the aid of fresh external heat, because sufficient heat[23] is evolved to raise the temperature of the remaining parts of the combustible substance to the required degree. Examples of this are familiar to all from every-day experience. Combustion proceeds in oxygen with greater rapidity, and is accompanied by a more powerful incandescence, than in ordinary air. This may be demonstrated by a number of very convincing experiments. If a piece of charcoal, attached to a wire and previously brought to red-heat, be plunged into a flask full of oxygen, it burns rapidly at a white heat--_i.e._ it combines with the oxygen, forming a gaseous product of combustion called carbonic anhydride, or carbonic acid gas, CO_{2}. This is the same gas that is evolved in the act of respiration, for charcoal is one of the substances which is obtained by the decomposition of all organic substances which contain it, and in the process of respiration part of the constituents of the body, so to speak, slowly burn. If a piece of burning sulphur be placed in a small cup attached to a wire and introduced into a flask full of oxygen, then the sulphur, which burns in air with a very feeble flame, burns in the oxygen with a violet flame, which, although pale, is much larger than in air. If the sulphur be exchanged for a piece of phosphorus,[24] then, unless the phosphorus be heated, it combines very slowly with the oxygen; but, if heated, although on only one spot, it burns with an exceedingly brilliant white flame. In order to heat the phosphorus inside the flask, the simplest way is to bring a red-hot wire into contact with it. Before the charcoal can burn, it must be brought to a state of incandescence. Sulphur also will not burn under 100°, whilst phosphorus inflames at 40°. Phosphorus which has been already lighted in air cannot so well be introduced into the flask, because it burns very rapidly and with a large flame in air. If a small lump of metallic _sodium_ be put in a small cup made of lime,[25] melted, and ignited,[26] it burns very feebly in air. But if burning sodium be introduced into oxygen, the combustion is invigorated and is accompanied by a brighter yellow flame. Metallic _magnesium_, which burns brightly in air, continues to burn with still greater vigour in oxygen, forming a white powder, which is a compound of magnesium with oxygen (magnesium oxide; magnesia). A strip of _iron_ or steel does not burn in air, but an iron wire or steel spring may be easily burnt in oxygen.[27] The combustion of steel or iron in oxygen is not accompanied by a flame, but sparks of oxide fly in all directions from the burning portions of the iron.[28]

[22] Certain substances (with which we shall afterwards become
acquainted), however, ignite spontaneously in air; for example,
impure phosphuretted hydrogen, silicon hydride, zinc ethyl, and
pyrophorus (very finely divided iron, &c.)

[23] If so little heat is evolved that the adjacent parts are not
heated to the temperature of combustion, then combustion will
cease.

[24] The phosphorus must be dry; it is usually kept in water, as
it oxidises in air. It should be cut under water, as otherwise
the freshly-cut surface oxidises. It must be dried carefully and
quickly by wrapping it in blotting-paper. If damp, it splutters
on burning. A small piece should be taken, as otherwise the iron
spoon will melt. In this and the other experiments on combustion,
water should be poured over the bottom of the vessel containing
the oxygen, to prevent it from cracking. The cork closing
the vessel should not fit tightly, in order to allow for the
expansion of the gas due to the heat of the combustion.

[25] An iron cup will melt with sodium in oxygen.

[26] In order to rapidly heat the lime crucible containing the sodium,
it is heated in the flame of a blowpipe described in Chap. VIII.

[27] In order to burn a watch spring, a piece of tinder (or paper
soaked in a solution of nitre, and dried) is attached to one
end. The tinder is lighted, and the spring is then plunged into
the oxygen. The burning tinder heats the end of the spring, the
heated part burns, and in so doing heats the further portions of
the spring, which then burns completely if sufficient oxygen be
present.

[28] The sparks of rust are produced, owing to the fact that the volume
of the oxide of iron is nearly twice that of the volume of the
iron, and as the heat evolved is not sufficient to entirely melt
the oxide or the iron, the particles must be torn off and fly
about. Similar sparks are formed in the combustion of iron, in
other cases also. We saw the combustion of iron filings in the
Introduction. In the welding of iron small iron splinters fly off
in all directions and burn in the air, as is seen from the fact
that whilst flying through the air they remain red hot, and also
because, on cooling, they are seen to be no longer iron, but a
compound of it with oxygen. The same thing takes place when the
hammer of a gun strikes against the flint. Small scales of steel
are heated by the friction, and glow and burn in the air. The
combustion of iron is still better seen by taking it as a very
fine powder, such as is obtained by the decomposition of certain
of its compounds--for instance, by heating Prussian blue, or by
the reduction of its compounds with oxygen by hydrogen; when this
fine powder is strewn in air, it burns by itself, even without
being previously heated (it forms a pyrophorus). This obviously
depends on the fact that the powder of iron presents a larger
surface of contact with air than an equal weight in a compact
form.

In order to demonstrate by experiment the _combustion of hydrogen_ in oxygen, a gas-conducting tube, bent so as to form a convenient jet, is led from the vessel evolving hydrogen. The hydrogen is first set light to in air, and then the gas-conducting tube is let down into a flask containing oxygen. The combustion in oxygen will be similar to that in air; the flame remains pale, notwithstanding the fact that its temperature rises considerably. It is instructive to remark that oxygen may burn in hydrogen, just as hydrogen in oxygen. In order to show the combustion of oxygen in hydrogen, a tube bent vertically upwards and ending in a fine orifice is attached to the stopcock of a gas-holder full of oxygen. Two wires, placed at such a distance from each other as to allow the passage of a constant series of sparks from a Ruhmkorff's coil, are fixed in front of the orifice of the tube. This is in order to ignite the oxygen, which may also be done by attaching tinder round the orifice, and burning it. When the wires are arranged at the orifice of the tube, and a series of sparks passes between them, then an inverted (because of the lightness of the hydrogen) jar full of hydrogen is placed over the gas-conducting tube. When the jar covers the orifice of the gas-conducting tube (and not before, as otherwise an explosion might take place) the cock of the gasometer is opened, and the oxygen flows into the hydrogen and is set light to by the sparks. The flame obtained is similar to that formed by the combustion of hydrogen in oxygen.[29] From this it is evident that the flame is the locality where the oxygen combines with the hydrogen, therefore a flame of burning oxygen can be obtained as well as a flame of burning hydrogen.

[29] The experiment may be conducted without the wires, if the hydrogen
be lighted in the orifice of an inverted cylinder, and at
the same time the cylinder be brought over the end of a
gas-conducting tube connected with a gas-holder containing
oxygen. Thomsen's method may be adopted for a lecture experiment.
Two glass tubes, with platinum ends, are passed through orifices,
about 1-1-1/2 centimetre apart, in a cork. One tube is connected
with a gas-holder containing oxygen, and the other with a
gas-holder full of hydrogen. Having turned on the gases, the
hydrogen is lighted, and a common lamp glass, tapering towards
the top, is placed over the cork. The hydrogen continues to burn
inside the lamp glass, at the expense of the oxygen. If the
current of oxygen be then decreased little by little, a point is
reached when, owing to the insufficient supply of oxygen, the
flame of the hydrogen increases in size, disappears for several
moments, and then reappears at the tube supplying the oxygen. If
the flow of oxygen be again increased, the flame reappears at the
hydrogen tube. Thus the flame may be made to appear at one or the
other tube at will, only the increase or decrease of the current
of gas must take place by degrees and not suddenly. Further, air
may be taken instead of oxygen, and ordinary coal-gas instead
of hydrogen, and it will then be shown how air burns in an
atmosphere of coal-gas, and it can easily be proved that the lamp
glass is full of a gas combustible in air, because it may be
lighted at the top.

If, instead of hydrogen, any other combustible gas be taken--for example, ordinary coal gas--then the phenomenon of combustion will be exactly the same, only a bright flame will be obtained, and the products of combustion will be different. However, as coal gas contains a considerable amount of free and combined hydrogen, it will also form a considerable quantity of water in its combustion.

If hydrogen be mixed with oxygen in the proportion in which they form water--_i.e._ if two volumes of hydrogen be taken for each volume of oxygen--then the mixture will be the same as that obtained by the decomposition of water by a galvanic current--detonating gas.

We have already mentioned in the last chapter that the combination of these gases, or their explosion, may be brought about by the action of an electric spark, because the spark heats the space through which it passes, and acts consequently in a manner similar to ignition by means of contact with an incandescent or burning substance.[29 bis] Cavendish made this experiment on the ignition of detonating gas, at the end of the last century, in the apparatus shown in fig. 31. Ignition by the aid of the electric spark is convenient, for the reason that it may then be brought about in a closed vessel, and hence chemists still employ this method when it is required to ignite a mixture of oxygen with a combustible gas in a closed vessel. For this purpose, especially since Bunsen's time,[30] an _eudiometer_ is employed. It consists of a thick glass tube graduated along its length in millimetres (for indicating the height of the mercury column), and calibrated for a definite volume (weight of mercury). Two platinum wires are fused into the upper closed end of the tube, as shown in fig. 32.[31] By the aid of the eudiometer we may not only determine the volumetric composition of water,[32] and the quantitative contents of oxygen in air,[33] but also make a number of experiments explaining the phenomenon of combustion.

[29 bis] In fact, instead of a spark a fine wire may be taken, and
an electric current passed through it to bring it to a state of
incandescence; in this case there will be no sparks, but the
gases will inflame if the wire be fine enough to become red hot
by the passage of the current.