Chapter III: Oxygen and the Chief Aspects of Its Saline Combinations (2)

[30] Now, a great many other different forms of apparatus, sometimes
designed for special purposes, are employed in the laboratory for
the investigation of gases. Detailed descriptions of the methods
of gas analysis, and of the apparatus employed, must be looked
for in works on analytical and applied chemistry.

[31] They must be sealed into the tube in such a manner as to leave
no aperture between them and the glass. In order to test this,
the eudiometer is filled with mercury, and its open end inverted
into mercury. If there be the smallest orifice at the wires, the
external air will enter into the cylinder and the mercury will
fall, although not rapidly if the orifice be very fine.

[32] The eudiometer is used for determining the composition of
combustible gases. A detailed account of _gas analysis_ would be
out of place in this work (_see_ Note 30), but, as an example,
we will give a short description of the determination of the
composition of water by the eudiometer.

Pure and dry oxygen is first introduced into the eudiometer. When
the eudiometer and the gas in it acquire the temperature of the
surrounding atmosphere--which is recognised by the fact of the
meniscus of the mercury not altering its position during a long
period of time--then the heights at which the mercury stands in
the eudiometer and in the bath are observed. The difference (in
millimetres) gives the height of the column of mercury in the
eudiometer. It must be reduced to the height at which the mercury
would stand at 0° and deducted from the atmospheric pressure, in
order to find the pressure under which the oxygen is measured
(_see_ Chap. I. Note 29). The height of the mercury also shows
the volume of the oxygen. The temperature of the surrounding
atmosphere and the height of the barometric column must also
be observed, in order to know the temperature of the oxygen
and the atmospheric pressure. When the volume of the oxygen
has been measured, pure and dry hydrogen is introduced into
the eudiometer, and the volume of the gases in the eudiometer
again measured. They are then exploded. This is done by a Leyden
jar, whose outer coating is connected by a chain with one wire,
so that a spark passes when the other wire, fused into the
eudiometer, is touched by the terminal of the jar. Or else an
electrophorus is used, or, better still, a Ruhmkorff's coil,
which has the advantage of working equally well in damp or dry
air, whilst a Leyden jar or electrical machine does not act in
damp weather. Further, it is necessary to close the lower orifice
of the eudiometer before the explosion (for this purpose the
eudiometer, which is fixed in a stand, is firmly pressed down
from above on to a piece of india-rubber placed at the bottom
of the bath), as otherwise the mercury and gas would be thrown
out of the apparatus by the explosion. It must also be remarked
that to ensure complete combustion the proportion between the
volumes of oxygen and hydrogen must not exceed twelve of hydrogen
to one volume of oxygen, or fifteen volumes of oxygen to one
volume of hydrogen, because no explosion will take place if one
of the gases be in great excess. It is best to take a mixture
of one volume of hydrogen with several volumes of oxygen. The
combustion will then be complete. It is evident that water is
formed, and that the volume (or tension) is diminished, so that
on opening the end of the eudiometer the mercury will rise in
it. But the tension of the aqueous vapour is now added to the
tension of the gas remaining after the explosion. This must be
taken into account (Chap. I. Note 1). If but little gas remain,
the water which is formed will be sufficient for its saturation
with aqueous vapour. This may be learnt from the fact that drops
of water are visible on the sides of the eudiometer after the
mercury has risen in it. If there be none, a certain quantity of
water must be introduced into the eudiometer. Then the number of
millimetres expressing the pressure of the vapour corresponding
with the temperature of the experiment must be subtracted from
the atmospheric pressure at which the remaining gas is measured,
otherwise the result will be inaccurate (Chap. I. Note 1).

This is essentially the method of the determination of the
composition of water which was made for the first time by
Gay-Lussac and Humboldt with sufficient accuracy. Their
determinations led them to the conclusion that water consists of
two volumes of hydrogen (more exactly 2·003, Le Duc 1892), and
one volume of oxygen. Every time they took a greater quantity of
oxygen, the gas remaining after the explosion was oxygen. When
they took an excess of hydrogen, the remaining gas was hydrogen;
and when the oxygen and hydrogen were taken in exactly the above
proportion, neither one nor the other remained. The composition
of water was thus definitely confirmed.

[33] Concerning this application of the eudiometer, see the chapter
on Nitrogen. It may be mentioned as illustrating the various uses
of the eudiometer that Prof. Timeraseeff employed microscopically
small eudiometers to analyse the bubbles of gas given off from
the leaves of plants.

Thus, for example, it may be demonstrated, by the aid of the eudiometer, that for the ignition of detonating gas, a _definite temperature_ is required. If the temperature be below that required, combination will not take place, but if at any spot within the tube it rises to the temperature of inflammation, then combination will ensue at that spot, and evolve enough heat for the ignition of the adjacent portions of the detonating mixture. If to 1 volume of detonating gas there be added 10 volumes of oxygen, or 4 volumes of hydrogen, or 3 volumes of carbonic anhydride, then we shall not obtain an explosion by passing a spark through the diluted mixture. This depends on the fact that the temperature falls with the dilution of the detonating gas by another gas, because the heat evolved by the combination of the small quantity of hydrogen and oxygen brought to incandescence by the spark is not only transmitted to the water proceeding from the combination, but also to the foreign substance mixed with the detonating gas.[34] The necessity of a definite temperature for the ignition of detonating gas is also seen from the fact that pure detonating gas explodes in the presence of a red-hot iron wire, or of charcoal heated to 275°, but with a lower degree of incandescence there is not any explosion. It may also be brought about by rapid compression, when, as is known, heat is evolved.[35] Experiments made in the eudiometer showed that the ignition of detonating gas takes place at a temperature between 450° and 560°.[36]

[34] Thus 1/4 volume of carbonic oxide, an equal volume of marsh gas,
two volumes of hydrogen chloride or of ammonia, and six volumes
of nitrogen or twelve volumes of air added to one volume of
detonating gas, prevent its explosion.

[35] If the compression be brought about slowly, so that the heat
evolved succeeds in passing to the surrounding space, then the
combination of the oxygen and hydrogen does not take place, even
when the mixture is compressed by 150 times; for the gases are
not heated. If paper soaked with a solution of platinum (in aqua
regia) and sal ammoniac be burnt, then the ash obtained contains
very finely-divided platinum, and in this form it is best fitted
for igniting hydrogen and detonating gas. Platinum wire requires
to be heated, but platinum in so finely divided a state as it
occurs in this ash inflames hydrogen, even at -20°. Many other
metals, such as palladium (175°), iridium, and gold, act with a
slight rise of temperature, like platinum; but mercury, at its
boiling point, does not inflame detonating gas, although the
slow formation of water then begins at 305°. All data of this
kind show that the explosion of detonating gas presents one of
the many cases of contact phenomena. This conclusion is further
confirmed by the researches of V. Meyer (1892). He showed that
only a very slow formation of steam begins at 448°, and that
it only proceeds more rapidly at 518°. The temperature of the
explosion of detonating gas, according to the same author,
varies according as to whether the explosion is produced in open
vessels or in closed tubes. In the first case the temperature
of explosion lies between 530°-606°, and in the second between
630°-730°. In general it may be remarked that the temperature of
explosion of gaseous mixtures is always lower in closed vessels
than when the detonating mixture flows freely through tubes.
According to Freyer and V. Meyer, the following gases when mixed
with the requisite amount of oxygen explode at the following
temperatures:

+----------------+---------------------+-------------------+
| -- | When flowing freely | In closed vessels |
+----------------+---------------------+-------------------+
| H_{2} | 630°-730° | 530°-606° |
| CH_{4} | 650°-730° | 606°-650° |
| C_{2}H_{6} | 606°-650° | 530°-606° |
| C_{2}H_{4} | 606°-650° | 530°-606° |
| CO | 650°-730° | 650°-730° |
| H_{2}S | 315°-320° | 250°-270° |
| H_{2} + Cl_{2} | 430°-440° | 240°-270° |
+----------------+---------------------+-------------------+

The velocity of the transmission of explosion in gaseous mixtures
is as characteristic a quantity for gaseous systems as the
velocity of the transmission of sound. Berthelot showed that this
velocity depends neither upon the pressure nor upon the size of
the tubes in which the gaseous mixture is contained, nor upon the
material out of which the tube is made. Dixon (1891) determined
the magnitude of these velocities for various mixtures, and his
results proved very near to those previously given by Berthelot.
For comparison we give the velocities expressed in metres per
second:

+-----------------+-------+-----------+
| -- | Dixon | Berthelot |
+-----------------+-------+-----------+
| H_{2} + O | 2,821 | 2,810 |
| H_{2} + N_{2}O | 2,305 | 2,284 |
| CH_{4} + 4O | 2,322 | 2,287 |
| C_{2}H_{2} + 6O | 2,364 | 2,210 |
| C_{2}H_{2} + 5O | 2,391 | 2,482 |
| C_{2}N_{2} + 4O | 2,321 | 2,195 |
+-----------------+-------+-----------+

The addition of oxygen to detonating gas lowers the velocity of
the transmission of explosion almost as much as the introduction
of nitrogen. An excess of hydrogen on the contrary raises the
velocity of transmission. It is remarked that the explosion
of mixtures of oxygen with marsh gas, ethylene and cyanogen
is transmitted more quickly if the oxygen be taken in such a
proportion that the carbon should burn to oxide of carbon, _i.e._
the velocity of the explosion is less if the oxygen be taken in
sufficient quantity to form carbonic anhydride. Observations upon
liquid and solid explosives (Berthelot) show that in this case
the velocity of transmission of explosion is dependent upon the
material of the tube. Thus the explosion of liquid nitro-methyl
ether in glass tubes travels at the rate (in dependence upon
the diam., from 1 mm.-45 mm.) of from 1,890 to 2,482 metres,
and in tubes of Britannia metal (3 mm. in diam) at the rate of
1,230 metres. The harder the tube the greater the velocity of
transmission of explosion. The following are the velocities for
certain bodies:

metres
Nitro-glycerine 1,300
Dynamite 2,500
Nitro-mannite 7,700
Picric acid 6,500

In conclusion we may add that Mallard and Le Chatelier (1882)
observed that in the explosion of a mixture of 1 volume of
detonating gas with _n_ volumes of an inert gas, the pressure is
approximately equal to 9·2-0·9_n_ atmospheres.

[36] From the very commencement of the promulgation of the idea of
dissociation, it might have been imagined that reversible
reactions of combination (the formation of H_{2} and O belongs
to this number) commence at the same temperature as that at
which dissociation begins. And in many cases this is so, but
not always, as may be seen from the facts (1) that at 450-560°,
when detonating gas explodes, the density of aqueous vapour not
only does not vary (and it hardly varies at higher temperatures,
probably because the amount of the products of dissociation is
small), but there are not, as far as is yet known, any traces
of dissociation; (2) that under the influence of contact the
temperature at which combination takes place falls even to the
ordinary temperature, when water and similar compounds naturally
are not dissociated and, judging from the data communicated
by D. P. Konovaloff (Introduction, Note 39) and others, it is
impossible to escape the phenomena of contact; all vessels,
whether of metal or glass, show the same influence as spongy
platinum, although to a much less degree. The phenomena of
contact, judging from a review of the data referring to it,
must be especially sensitive in reactions which are powerfully
exothermal, and the explosion of detonating gas is of this kind.

The combination of hydrogen with oxygen is accompanied by the evolution of a very considerable amount of heat; according to the determinations of _Favre_ and _Silbermann_,[37] 1 part by weight of hydrogen in forming water evolves 34,462 units of heat. Many of the most recent determinations are very close to this figure, so that it may be taken that in the formation of 18 parts of water (H_{2}O) there are evolved 69 major calories, or 69,000 units of heat.[38] _If the specific heat of aqueous vapour_ (0·48) _remained constant from the ordinary temperature to that at which the combustion of detonating gas takes place_ (but there is now no doubt that it increases), were the combustion concentrated at one point[39] (but it occurs in the whole region of a flame), were there no loss from radiation and heat conduction, and _did dissociation not take place_--that is, did not a state of equilibrium between the hydrogen, oxygen, and water come about--_then it would be possible to calculate the temperature of the flame of detonating gas_. It would then be 8,000°.[40] In reality it is very much lower, but it is nevertheless higher than the temperature attained in furnaces and flames, and is as high as 2,000°. The explosion of detonating gas is explained by this high temperature, because the aqueous vapour formed must occupy a volume at least 5 times greater than that occupied by the detonating gas at the ordinary temperature. Detonating gas emits a sound, not only as a consequence of the commotion which occurs from the rapid expansion of the heated vapour, but also because it is immediately followed by a cooling effect, the conversion of the vapour into water, and a rapid contraction.[41]

[37]

The amount of heat evolved in the combustion of a known weight
(for instance, 1 gram) of a given substance is determined by the
rise in temperature of water, to which the whole of the heat
evolved in the combustion is transmitted. A _calorimeter_, for
example that shown in fig. 33, is employed for this purpose.
It consists of a thin (in order that it may absorb less heat),
polished (that it should transmit a minimum of heat) metallic
vessel, surrounded by down (_c_), or some other bad conductor
of heat, and an outer metallic vessel. This is necessary in
order that the least possible amount of heat should be lost
from the vessels; nevertheless, there is always a certain loss,
whose magnitude is determined by preliminary experiment (by
taking warm water, and determining its fall in temperature after
a definite period of time) as a correction for the results
of observations. The water to which the heat of the burning
substance is transmitted is poured into the vessel. The stirrer
_g_ allows of all the layers of water being brought to the same
temperature, and the thermometer serves for the determination
of the temperature of the water. The heat evolved passes,
naturally, not to the water only, but to all the parts of the
apparatus. The quantity of water corresponding to the whole
amount of those objects (the vessels, tubes, &c.) to which the
heat is transmitted is previously determined, and in this manner
another most important correction is made in the calorimetric
determinations. The combustion itself is carried on in the vessel
_a_. The ignited substance is introduced through the tube at the
top, which closes tightly. In fig. 33 the apparatus is arranged
for the combustion of a gas, introduced by a tube. The oxygen
required for the combustion is led into _a_ by the tube _e_, and
the products of combustion either remain in the vessel _a_ (if
liquid or solid), or escape by the tube _f_ into an apparatus in
which their quantity and properties can easily be determined.
Thus the heat evolved in combustion passes to the walls of the
vessel _a_, and to the gases which are formed in it, and these
transmit it to the water of the calorimeter.

[38] This quantity of heat corresponds with the formation of
liquid water at the ordinary temperature from detonating gas at
the same temperature. If the water be as vapour the heat evolved
= 58 major calories; if as ice = 70·4 major calories. A portion
of this heat is due to the fact that 2 vols. of hydrogen and 1
vol. of oxygen give 2 vols. of aqueous vapour--that is to say,
contraction ensues--and this evolves heat. This quantity of heat
may be calculated, but it cannot be said how much is expended
in the separation of the atoms of oxygen from each other, and,
therefore, strictly speaking, we do not know the quantity of
heat which is evolved in the reaction alone, although the number
of units of heat evolved in the combustion of detonating gas is
accurately known.

The construction of the calorimeter and even the method of
determination vary considerably in different cases. Since the
beginning of the nineties, a large number of determinations
of the heat of combustion have been conducted in closed bombs
containing compressed oxygen. The greatest number of calorimetric
determinations were made by Berthelot and Thomsen. They are
given in their works _Essai de mécanique chimique fondée
sur la thermochimie_, by M. Berthelot, 1879 (2 vols.), and
_thermochemische Untersuchungen_, by J. Thomsen, 1886 (4 vols.)
The most important methods of recent thermochemistry, and all
the trustworthy results of experiment, are given in Prof. P. F.
Louginin's _Description of the Different Modes of Determining
the Heat of Combustion of Organic Compounds_, Moscow, 1894.
The student must refer to works on theoretical and physical
chemistry for a description of the elements and methods of
_thermochemistry_, into the details of which it is impossible to
enter in this work. One of the originators of thermochemistry,
Hess, was a member of the St. Petersburg Academy of Sciences.
Since 1870 a large amount of research has been carried out in
this province of chemistry, especially in France and Germany,
after the investigations of the French Academician, Berthelot,
and Professor Thomsen, of Copenhagen. Among Russians, Beketoff,
Louginin, Cheltzoff, Chroustchoff, and others are known by their
thermochemical researches. The present epoch of thermochemistry
must be considered rather as a collective one, wherein the
material of facts is amassed, and the first consequences arising
from them are noticed. In my opinion two essential circumstances
prevent the possibility of deducing any exact consequences, of
importance to chemical mechanics, from the immense store of
thermochemical data already collected: (1) The majority of the
determinations are conducted in weak aqueous solutions, and, the
heat of solution being known, are referred to the substances
in solution; yet there is much (Chapter I.) which leads to the
conclusion that in solution water does not play the simple
part of a diluting medium, but of itself acts independently
in a chemical sense on the substance dissolved. (2) Physical
and mechanical changes (decrease of volume, diffusion, and
others) invariably proceed side by side with chemical changes,
and for the present it is impossible, in a number of cases, to
distinguish the thermal effect of the one and the other kind of
change. It is evident that the one kind of change (chemical) is
essentially inseparable and incomprehensible without the other
(mechanical and physical); and therefore it seems to me that
thermochemical data will only acquire their true meaning when
the connection between the phenomena of both kinds (on the one
hand chemical and atomic, and on the other hand mechanical and
molecular or between entire masses) is explained more clearly
and fully than is at present the case. As there is no doubt that
the simple mechanical contact, or the action of heat alone,
on substances sometimes causes an evident and always a latent
(incipient) chemical change--that is, a different distribution
or motion of the atoms in the molecules--it follows that purely
chemical phenomena are inseparable from physical and mechanical
phenomena. A mechanical change may be imagined without a physical
change, and a physical without a chemical change, but it is
impossible to imagine a chemical change without a physical and
mechanical one, for without the latter we should not be able
to recognise the former, and it is by their means that we are
enabled to do so.

[39] The flame, or locality where the combustion of gases and vapours
takes place, is a complex phenomenon, 'an entire factory,' as
Faraday says, and therefore we will consider flame in some detail
in one of the following notes.

[40] If 34,500 units of heat are evolved in the combustion of 1 part
of hydrogen, and this heat is transmitted to the resulting 9
parts by weight of aqueous vapour, then we find that, taking the
specific heat of the latter as 0·475, each unit of heat raises
the temperature of 1 part by weight of aqueous vapour 2°·1 and 9
parts by weight (2·1 ÷ 9) O°·23; hence the 34,500 units of heat
raise its temperature 7,935°. If detonating gas is converted
into water in a closed space, then the aqueous vapour formed
cannot expand, and therefore, in calculating the temperature of
combustion, the specific heat at a constant volume must be taken
into consideration; for aqueous vapour it is 0·36. This figure
gives a still higher temperature for the flame. In reality it
is much lower, but the results given by different observers are
very contradictory (from 1,700° to 2,400°), the discrepancies
depending on the fact that flames of different sizes are cooled
by radiation to a different degree, but mainly on the fact that
the methods and apparatus (pyrometers) for the determination
of high temperatures, although they enable relative changes of
temperature to be judged, are of little use for determining their
absolute magnitude. By taking the temperature of the flame of
detonating gas as 2,000°, I give, I think, the average of the
most trustworthy determinations and calculations based upon the
determination of the variation of the specific heat of aqueous
vapour and other gases (_see_ Chapter XLI.)

[41] It is evident that not only hydrogen, but every other combustible
gas, will give an explosive mixture with oxygen. For this reason
coal-gas mixed with air explodes when the mixture is ignited.
The pressure obtained in the explosions serves as the _motive
power of gas engines_. In this case advantage is taken, not only
of the pressure produced by the explosion, but also of that
contraction which takes place after the explosion. On this is
based the construction of several motors, of which Lenoir's was
formerly, and Otto's is now, the best known. The explosion is
usually produced by coal-gas and air, but of late the vapours of
combustible liquids (kerosene, benzene) are also being employed
in place of gas (Chapter IX.) In Lenoir's engine a mixture of
coal-gas and air is ignited by means of sparks from a Ruhmkorff's
coil, but in the most recent machines the gases are ignited by
the direct action of a gas jet, or by contact with the hot walls
of a side tube.

Mixtures of hydrogen and of various other gases with oxygen are taken advantage of for obtaining high temperatures. By the aid of such high temperatures metals like platinum may be melted on a large scale, which cannot be performed in furnaces heated with charcoal and fed by a current of air. The burner, shown in fig. 34, is constructed for the application of detonating gas to the purpose. It consists of two brass tubes, one fixed inside the other, as shown in the drawing. The internal central tube C C conducts oxygen, and the outside, enveloping, tube E' E' conducts hydrogen. Previous to their egress the gases do not mix together, so that there can be no explosion inside the apparatus. When this burner is in use C is connected with a gas-holder containing oxygen, and E with a gas-holder containing hydrogen (or sometimes coal-gas). The flow of the gases can be easily regulated by the stopcocks O H. The flame is shortest and evolves the greatest heat when the gases burning are in the proportion of 1 volume of oxygen to 2 volumes of hydrogen. The degree of heat may be easily judged from the fact that a thin platinum wire placed in the flame of a properly proportioned mixture easily melts. By placing the burner in the orifice of a hollow piece of lime, a crucible A B is obtained in which the platinum may be easily melted, even in large quantities if the current of oxygen and hydrogen be sufficiently great (Deville). The flame of detonating gas may also be used for illuminating purposes. It is by itself very pale, but owing to its high temperature it may serve for rendering infusible objects incandescent, and at the very high temperature produced by the detonating gas the incandescent substance gives a most intense light. For this purpose lime, magnesia, or oxide of zirconium are used, as they are not fusible at the very high temperature evolved by the detonating gas. A small cylinder of lime placed in the flame of detonating gas, if regulated to the required point, gives a very brilliant white light, which was at one time proposed for illuminating lighthouses. At present in the majority of cases the electric light, owing to its constancy and other advantages, has replaced it for this purpose. The light produced by the incandescence of lime in detonating gas is called the _Drummond light_ or _limelight_.

The above cases form examples of the combustion of elements in oxygen, but exactly similar phenomena are observed in the _combustion of compounds_. So, for instance, the solid, colourless, shiny substance, naphthalene, C_{10}H_{8}, burns in the air with a smoky flame, whilst in oxygen it continues to burn with a very brilliant flame. Alcohol, oil, and other substances burn brilliantly in oxygen on conducting the oxygen by a tube to the flame of lamps burning these substances. A high temperature is thus evolved, which is sometimes taken advantage of in chemical practice.

In order to understand why combustion in oxygen proceeds more rapidly, and is accompanied by a more intense heat effect, than combustion in air, it must be recollected that air is oxygen diluted with nitrogen, which does not support combustion, and therefore fewer particles of oxygen flow to the surface of a substance burning in air than when burning in pure oxygen, besides which the reason of the intensity of combustion in oxygen is the high temperature acquired by the substance burning in it.[41 bis]

[41 bis] Let us consider as an example the combustion of
sulphur in air and in oxygen. If 1 gram of sulphur burns in air
or oxygen it evolves in either case 2250 units of heat--_i.e._
evolves sufficient heat for heating 2,250 grams of water 1°
C. This heat is first of all transmitted to the sulphurous
anhydride, SO_{2}, formed by the combination of sulphur with
oxygen. In its combustion 1 gram of sulphur forms 2 grams of
sulphurous anhydride--_i.e._ the sulphur combines with 1 gram of
oxygen. In order that 1 gram of sulphur should have access to 1
gram of oxygen in air, it is necessary that 3·4 grams of nitrogen
should simultaneously reach the sulphur, because air contains
seventy-seven parts of nitrogen (by weight) per twenty-three
parts of oxygen. Thus in the combustion of 1 gram of sulphur,
the 2,250 units of heat are transmitted to 2 grams of sulphurous
oxide and to at least 3·4 grams of nitrogen. As 0·155 unit of
heat is required to raise 1 gram of sulphurous anhydride 1°
C., therefore 2 grams require 0·31 unit. So also 3·4 grams of
nitrogen require 3·4 × 0·244 or 0·83 unit of heat, and therefore
in order to raise both gases 1° C. 0·31 + 0·83 or 1·14 unit of
heat is required; but as the combustion of the sulphur evolves
2,250 units of heat, therefore the gases might be heated (if
their specific heats remained constant) to 2250/1·14 or 1,974°
C. That is, the maximum possible temperature of the flame of the
sulphur burning in air will be 1,974° C. In the combustion of the
sulphur in oxygen the heat evolved (2,250 units) can only pass to
the 2 grams of sulphurous anhydride, and therefore the highest
possible temperature of the flame of the sulphur in oxygen will
be = 2250/0·31 or 7258°. In the same manner it may be calculated
that the temperature of charcoal burning in air cannot exceed
2,700°, while in oxygen it may attain 10,100° C. For this reason
the temperature in oxygen will always be higher than in air,
although (judging from what has been said respecting detonating
gas) neither one temperature nor the other will ever approximate
to the theoretical amount.

Among the phenomena accompanying the combustion of certain substances, the _phenomenon of flame_ attracts attention. Sulphur, phosphorus, sodium, magnesium, naphthalene, &c., burn like hydrogen with a flame, whilst in the combustion of other substances no flame is observed, as, for instance, in the combustion of iron and of charcoal. The appearance of flame depends on the capacity of the combustible substance to yield gases or vapours at the temperature of combustion. At the temperature of combustion, sulphur, phosphorus, sodium, and naphthalene pass into vapour, whilst wood, alcohol, oil, &c., are decomposed into gaseous and vaporous substances. The combustion of gases and vapours forms flames, and therefore _a flame is composed of the hot and incandescent gases and vapours produced by combustion_. It may easily be proved that the flames of such non-volatile substances as wood contain volatile and combustible substances formed from them, by placing a tube in the flame connected with an aspirator. Besides the products of combustion, combustible gases and liquids, previously in the flame as vapours, collect in the aspirator. For this experiment to succeed--_i.e._ in order to really extract combustible gases and vapours from the flame it is necessary that the suction tube should be placed _inside_ the flame. The combustible gases and vapours can only remain unburnt inside the flame, for at the surface of the flame they come into contact with the oxygen of the air and burn.[42] Flames are of different degrees of _brilliancy_, according to whether _solid_ incandescent particles occur in the combustible gas or vapour, or not. Incandescent gases and vapours emit but little light by themselves, and therefore give a paler flame.[43] If a flame does not contain solid particles it is transparent, pale, and emits but little light.[44] The flames of burning alcohol, sulphur, and hydrogen are of this kind. A pale flame may be rendered luminous by placing fine particles of solid matter in it. Thus, if a very fine platinum wire be placed in the pale flame of burning alcohol--or, better still, of hydrogen--the flame emits a bright light. This is still better seen by sifting the powder of an incombustible substance, such as fine sand, into the flame, or by placing a bunch of asbestos threads in it. Every brilliant flame always contains some kind of solid particles, or at least some very dense vapour. The flame of sodium burning in oxygen has a brilliant yellow colour, from the presence of particles of solid sodium oxide. The flame of magnesium is brilliant from the fact that in burning it forms solid magnesia, which becomes white hot, and similarly the brilliancy of the Drummond light is due to the heat of the flame raising the solid non-volatile lime to a state of incandescence. The flames of a candle, wood, and similar substances are brilliant, because they contain particles of charcoal or soot. It is not the flame itself which is luminous, but the incandescent soot it contains. These particles of charcoal which occur in flames may be easily observed by introducing a cold object, like a knife, into the flame.[45] The particles of charcoal burn at the outer surface of the flame if the supply of air be sufficient, but if the supply of air--that is, of oxygen--be insufficient for their combustion the flame smokes, because the unconsumed particles of charcoal are carried off by the current of air.[46]

[42] Faraday proved this by a very convincing experiment on
a candle flame. If one arm of a bent glass tube be placed in a
candle flame above the wick in the dark portion of the flame,
then the products of the partial combustion of the stearin will
pass up the tube, condense in the other arm, and collect in a
flask placed under it (fig. 35) as heavy white fumes which burn
when lighted. If the tube be raised into the upper luminous
portion of the flame, then a dense black smoke which will not
inflame accumulates in the flask. Lastly, if the tube be let down
until it touches the wick, then little but stearic acid condenses
in the flask.

[43] All transparent substances which transmit light with great ease
(that is, which absorb but little light) are but little luminous
when heated; so also substances which absorb but few heat rays,
when heated transmit few rays of heat.

[44] There is, however, no doubt but that very heavy dense vapours or
gases under pressure (according to the experiments of Frankland)
are luminous when heated, because, as they become denser they
approach a liquid or solid state. Thus detonating gas when
exploded under pressure gives a brilliant light.

[45] If hydrogen gas be passed through a volatile liquid hydrocarbon--for
instance, through benzene (the benzene may be poured directly
into the vessel in which hydrogen is generated)--then its vapour
burns with the hydrogen and gives a very bright flame, because
the resultant particles of carbon (soot) become incandescent.
Benzene, or platinum gauze, introduced into a hydrogen flame may
be employed for illuminating purposes.

[46] In _flames_ the separate parts may be distinguished with more or
less distinctness. That portion of the flame whither the
combustible vapours or gases flow, is not luminous because its
temperature is still too low for the process of combustion to
take place in it. This is the space which in a candle surrounds
the wick, or in a gas jet is immediately above the orifice from
which the gas escapes. In a candle the combustible vapours and
gases which are formed by the action of heat on the melted
tallow or stearin rise in the wick, and are heated by the high
temperature of the flame. By the action of the heat, the solid or
liquid substance is here, as in other cases, decomposed, forming
products of dry distillation. These products occur in the central
portion of the flame of a candle. The air travels to it from the
outside, and is not able to intermix at once with the vapours
and gases in all parts of the flame equally; consequently, in
the outer portion of the flame the amount of oxygen will be
greater than in the interior portions. But, owing to diffusion,
the oxygen, of course mixed with nitrogen, flowing towards the
combustible substance, does finally penetrate to the interior
of the flame (when the combustion takes place in ordinary air).
The combustible vapours and gases combine with this oxygen,
evolve a considerable amount of heat, and bring about that state
of incandescence which is so necessary both for keeping up the
combustion and also for the uses to which the flame is applied.
Passing from the colder envelope of air through the interior
of the flame, to the source of the combustible vapours (for
instance, the wick), we evidently first traverse layers of higher
and higher temperature, and then portions which are less and
less hot, in which the combustion is less complete, owing to the
limited supply of oxygen.

Thus unburnt products of the decomposition of organic substances
occur in the interior of the flame. But there is always free
hydrogen in the interior of the flame, even when oxygen is
introduced there, or when a mixture of hydrogen and oxygen burns,
because the temperature evolved in the combustion of hydrogen
or the carbon of organic matter is so high that the products
of combustion are themselves partially decomposed--that is,
dissociated--at this temperature. Hence, in a flame a portion
of the hydrogen and of the oxygen which might combine with
the combustible substances must always be present in a free
state. If a hydrocarbon burns, and we imagine that a portion
of the hydrogen is in a free state, then a portion of the
carbon must also occur in the same form in the flame, because,
other conditions being unchanged, carbon burns after hydrogen,
and this is actually observed in the combustion of various
hydrocarbons. Charcoal, or the soot of a common flame, arises
from the dissociation of organic substances contained in the
flame. The majority of hydrocarbons, especially those containing
much carbon--for instance, naphthalene--burn, even in oxygen,
with separation of soot. In that portion of the flame where the
hydrogen burns the carbon remains unburnt, or at least partly so.
It is this free carbon which causes the brilliancy of the flame.
That the interior of the flame contains a mixture which is still
capable of combustion may be proved by the following experiment:
A portion of the gases may be withdrawn by an aspirator from
the central portion of the flame of carbonic oxide, which is
combustible in air. For this purpose Deville passed water
through a metallic tube having a fine lateral orifice, which is
placed in the flame. As the water flows along the tube portions
of the gases of the flame enter, and, passing along the tube
alternately with cylinders of water, are carried away into an
apparatus where they can be investigated. It appears that all
portions of the flame obtained by the combustion of a mixture
of carbonic oxide and oxygen contain a portion of this mixture
still unburnt. The researches of Deville and Bunsen showed that
in the explosion of a mixture of hydrogen and of carbonic oxide
with oxygen in a closed space, complete combustion does not ever
take place immediately. If two volumes of hydrogen and one volume
of oxygen be confined in a closed space, then on explosion the
pressure does not attain that magnitude which it would were there
immediate and complete combustion. It may be calculated that the
pressure should attain twenty-six atmospheres. In reality, it
does not exceed nine and a half atmospheres.

Hence the admixture of the products of combustion with an
explosive mixture prevents the combustion of the remaining
mass, although capable of burning. The admixture of carbonic
anhydride prevents carbonic oxide from burning. The presence of
any other foreign gas interferes in the same manner. This shows
that every portion of a flame must contain combustible, burning,
and already burnt substances--_i.e._ oxygen, carbon, carbonic
oxide, hydrogen, hydrocarbons, carbonic anhydride, and water.
Consequently, _it is impossible to attain instantaneous complete
combustion_, and this is one of the reasons of the phenomenon
of flame. A certain space is required, and the temperature must
be unequal in different parts of it. In this space different
quantities of the component parts are successively subjected
to combustion, or are cooled under the influence of adjacent
objects, and combustion only ends where the flame ends. If
the combustion could be concentrated at one spot, then the
temperature would be incomparably higher than it is under the
actual circumstances.

The various regions of the flame have formed the frequent
subject of experimental research, and the experiments conducted
by Smithells and Ingle (1892) are particularly instructive;
they show that the reducing (interior) and oxidising (exterior)
portions of the flame of a burning gas may be divided by taking a
Bunsen burner and surrounding the flame of the gas burnt in it,
by another wider tube (without the access of air to the annular
space or allowing only a small current of air to pass), when
a gaseous mixture, containing oxide of carbon and capable of
further combustion, will issue from this enveloping tube, so that
a second flame, corresponding to the exterior (oxidising) portion
of an ordinary flame, may be obtained above the enveloping tube.
This division of the flame into two portions is particularly
clear when cyanogen C_{2}N_{2} is burnt, because the interior
portion (where CO is chiefly formed according to the equation
C_{2}N_{2} + O_{2} = 2CO + N_{2}, but a portion of the nitrogen
is oxidised) is of a rose colour, while the exterior portion
(where the CO burns into CO_{2} at the expense of a fresh
quantity of oxygen and of the oxides of nitrogen proceeding from
the interior portions) is of a bluish-grey colour.

The combination of various substances with oxygen may not present any signs of combustion--that is, the temperature may rise but inconsiderably. This may either proceed from the fact that the reaction of the substance (for example, tin, mercury, lead at a high temperature, or a mixture of pyrogallol with caustic potash at the ordinary temperature) evolves but little heat, or that the heat evolved is transmitted to good conductors of heat, like metals, or that the combination with oxygen takes place so slowly that the heat evolved succeeds in passing to the surrounding objects. Combustion is only a particular, intense, and evident case of combination with oxygen. Respiration is also an act of combination with oxygen; it also serves, like combustion, for the development of heat by those chemical processes which accompany it (the transformation of oxygen into carbonic anhydride). Lavoisier enunciated this in the lucid expression, 'respiration is slow combustion.'

Reactions involving slow combination of substances with oxygen are termed _oxidations_. Combination of this kind (and also combustion) often results in the formation of acid substances, and hence the name _oxygen_ (_Sauerstoff_). Combustion is only rapid oxidation. Phosphorus, iron, and wine may be taken as examples of substances which slowly oxidise in air at the ordinary temperature. If such a substance be left in contact with a definite volume of air or oxygen, it absorbs the oxygen little by little, as may be seen by the decrease in volume of the gas. This slow oxidation is not often accompanied by a sensible evolution of heat; an evolution of heat really does occur, only it is not apparent to our senses owing to the small rise in temperature which takes place; this is owing to the slow rate of the reaction and to the transmission of the heat formed as radiant heat, &c. Thus, in the oxidation of wine and its transformation into vinegar by the usual method of preparation of the latter, the heat evolved cannot be observed because it extends over several weeks, but in the so-called rapid process of the manufacture of vinegar, when a large quantity of wine is comparatively rapidly oxidised, the evolution of heat is quite apparent.

Such slow processes of oxidation are always taking place in nature by the action of the atmosphere. Dead organisms and the substances obtained from them--such as bodies of animals, wood, wool, grass, &c.--are especially subject to this action. They _rot_ and _decompose_--that is, their solid matter is transformed into gases, under the influence of moisture and atmospheric oxygen, and generally under the influence of other organisms, such as moulds, worms, micro-organisms (bacteria), and the like. These are processes of slow combustion, of slow combination with oxygen. It is well known that manure rots and develops heat, that stacks of damp hay, damp flour, straw, &c., become heated and are changed in the process.[47] In all these transformations the same chief products of combustion are formed as those which are contained in smoke; the carbon gives carbonic anhydride, and the hydrogen water. Hence these processes require oxygen just like combustion. This is the reason why the entire prevention of access of air hinders these transformations,[48] and an increased supply of air accelerates them. The mechanical treatment of arable lands by the plough, harrow, and other similar means has not only the object of facilitating the spread of roots in the ground, and of making the soil more permeable to water, but it also serves to facilitate the access of the air to the component parts of the soil; as a consequence of which the organic remains of soil rot--so to speak, breathe air and evolve carbonic anhydride. One acre of good garden land in the course of a summer evolves more than sixteen tons of carbonic anhydride.

[47] Cotton waste (used in factories for cleaning machines from
lubricating oil) soaked in oil and lying in heaps is
self-combustible, being oxidised by the air.

[48] When it is desired to preserve a supply of vegetable and animal
food, the access of the oxygen of the atmosphere (and also of
the germs of organisms present in the air) is often prevented.
With this object articles of food are often kept in hermetically
closed vessels, from which the air has been withdrawn; vegetables
are dried and soldered up while hot in tin boxes; sardines are
immersed in oil, &c. The removal of water from substances is also
sometimes resorted to with the same object (the drying of hay,
corn, fruits), as also is saturation with substances which absorb
oxygen (such as sulphurous anhydride), or which hinder the growth
of organisms forming the first cause of putrefaction, as in
processes of smoking, embalming, and in the keeping of fishes and
other animal specimens in spirit, &c.

It is not only vegetable and animal substances which are subject to slow oxidation in the presence of water. Some metals even rust under these conditions. Copper very easily absorbs oxygen in the presence of acids. Many metallic sulphides (for example, pyrites) are very easily oxidised with access of air and moisture. Thus processes of slow oxidation proceed throughout nature. However, there are many elements which do not under any circumstances combine directly with gaseous oxygen; nevertheless their compounds with oxygen may be obtained. Platinum, gold, iridium, chlorine, and iodine are examples of such elements. In this case recourse is had to a so-called _indirect method_--_i.e._ the given substance is combined with another element, and by a method of double decomposition this element is replaced by oxygen. Substances which do not directly combine with oxygen, but form compounds with it by an indirect method, often readily lose the oxygen which they had absorbed by double decomposition or at the moment of its evolution. Such, for example, are the compounds of oxygen with chlorine, nitrogen, and platinum, which evolve oxygen on heating--that is, they may be used as oxidising agents. In this respect _oxidising agents_, or those compounds of oxygen which are employed in chemical and technical practice for transferring oxygen to other substances, are especially remarkable. The most important among these is nitric acid or _aqua fortis_--a substance rich in oxygen, and capable of evolving it when heated, which easily oxidises a great number of substances. Thus nearly all metals and organic substances containing carbon and hydrogen are more or less oxidised when heated with nitric acid. If strong nitric acid be taken, and a piece of burning charcoal be immersed in the acid, it continues to burn. Chromic acid acts like nitric acid; alcohol burns when mixed with it. Although the action is not so marked, even water may oxidise with its oxygen. Sodium is not oxidised in perfectly dry oxygen at the ordinary temperature, but it burns very easily in water and aqueous vapour. Charcoal can burn in carbonic anhydride--a product of combustion--forming carbonic oxide. Magnesium burns in the same gas, separating carbon from it. Speaking generally, combined oxygen can pass from one compound to another.

The products of combustion or oxidation--and in general the definite compounds of oxygen--are termed _oxides_. Some oxides are not capable of combining with other oxides--or combine with only a few, and then with the evolution of very little heat; others, on the contrary, enter into combination with very many other oxides, and in general have remarkable chemical energy. The oxides incapable of combining with others, or only showing this quality in a small degree, are termed _indifferent oxides_. Such are the peroxides, of which mention has before been made.

The class of oxides capable of entering into mutual combination we will term _saline oxides_. They fall into two chief groups--at least, as regards the most extreme members. The members of one group combine with the members of the other group with particular ease. As representative of one group may be taken the oxides of the metals, magnesium, sodium, calcium, &c. Representatives of the other group are the oxides formed by the non-metals, sulphur, phosphorus, carbon. Thus, if we take the oxide of calcium, or lime, and bring it into contact with oxides of the second group, combination very readily ensues. For instance, if we mix calcium oxide with oxide of phosphorus they combine with great facility and with the evolution of much heat. If we pass the vapour of sulphuric anhydride, obtained by the combination of sulphurous oxide with oxygen, over pieces of lime heated to redness, the sulphuric anhydride is absorbed by the lime with the formation of a substance called calcium sulphate. The oxides of the first kind, which contain metals, are termed _basic oxides_ or _bases_. Lime is a familiar example of this class. The oxides of the second group, which are capable of combining with the bases, are termed _anhydrides of the acids_ or _acid oxides_. Sulphuric anhydride, SO_{3}, may be taken as a type of the latter group. It is a compound of sulphur with oxygen formed not directly but by the addition of a fresh quantity of oxygen to sulphurous anhydride, SO_{2}, by passing it together with oxygen over incandescent spongy platinum. Carbonic anhydride (often termed 'carbonic acid'), CO_{2}, phosphoric anhydride, sulphurous anhydride, are all acid oxides, for they can combine with such oxides as lime or calcium oxide, magnesia or magnesium oxide, MgO, soda or sodium oxide, Na_{2}O, &c.

If a given element form but one basic oxide, it is termed the _oxide_; for example, calcium oxide, magnesium oxide, potassium oxide. Some indifferent oxides are also called 'oxides' if they have not the properties of peroxides, and at the same time do not show the properties of acid anhydrides--for example, carbonic oxide, of which mention has already been made. If an element forms two basic oxides (or two indifferent oxides not having the characteristics of a peroxide) then that of the lower degree of oxidation is called a _suboxide_--that is, suboxides contain less oxygen than oxides. Thus, when copper is heated to redness in a furnace it increases in weight and absorbs oxygen, until for 63 parts of copper there is absorbed not more than 8 parts of oxygen by weight, forming a red mass, which is suboxide of copper; but if the roasting be prolonged, and the draught of air increased, 63 parts of copper absorb 16 parts of oxygen, and form black oxide of copper. Sometimes to distinguish between the degrees of oxidation a change of suffix is made in the oxidised element, _-ic_ oxide denoting the higher degree of oxidation, and _-ous_ oxide the lower degree. Thus ferrous oxide and ferric oxide are the same as suboxide of iron and oxide of iron. If an element forms one anhydride only, then it is named by an adjective formed from the name of the element made to end in _-ic_ and the word _anhydride_. When an element forms two anhydrides, then the suffixes _-ous_ and _-ic_ are used to distinguish them: _-ous_ signifying less oxygen than _-ic_; for example, sulphurous and sulphuric anhydrides.[49] When several oxides are formed from the same element, the prefixes _mon_, _di_, _tri_, _tetra_ are used, thus: chlorine monoxide, chlorine dioxide, chlorine trioxide, and chlorine tetroxide or chloric anhydride.