THE COMPOSITION OF WATER, HYDROGEN
The question now arises, Is not _water_ itself a _compound substance_?
Cannot it be formed by the mutual combination of some component parts?
Cannot it be broken up into its component parts? There cannot be the
least doubt that if it does split up, and if it is a compound, then it
is a _definite_ one characterised by the stability of the union between
those component parts from which it is formed. From the fact alone that
water passes into all physical states as a homogeneous whole, without
in the least varying chemically in its properties and without splitting
up into its component parts (neither solutions nor many hydrates can
be distilled--they are split up), we must conclude, from this fact
alone, that if water is a compound then it is a stable and definite
chemical compound capable of entering into many other combinations.
Like many other great discoveries in the province of chemistry, it is
to the end of the last century that we are indebted for the important
discovery that water is not a simple substance, that it is composed of
two substances like a number of other compound substances. This was
proved by two of the methods by which the compound nature of bodies
may be directly determined; by analysis and by synthesis--that is,
by a method of the decomposition of water into, and of the formation
of water from, its component parts. In 1781 Cavendish first obtained
water by burning hydrogen in oxygen, both of which gases were already
known to him. He concluded from this that water was composed of two
substances. But he did not make more accurate experiments, which would
have shown the relative quantities of the component parts in water,
and which would have determined its complex nature with certainty.
Although his experiments were the first, and although the conclusion
he drew from them was true, yet such novel ideas as the complex nature
of water are not easily recognised so long as there is no series of
researches which entirely and indubitably proves the truth of such a
conclusion. The fundamental experiments which proved the complexity
of water by the method of synthesis, and of its formation from other
substances, were made in 1789 by Monge, Lavoisier, Fourcroy, and
Vauquelin. They obtained four ounces of water by burning hydrogen,
and found that water consists of 15 parts of hydrogen and 85 parts of
oxygen. It was also proved that the weight of water formed was equal
to the sum of the weights of the component parts entering into its
composition; consequently, water contains all the matter entering into
oxygen and hydrogen. The complexity of water was proved in this manner
by a method of synthesis. But we will turn to its analysis--_i.e._ to
its decomposition into its component parts. The analysis may be more or
less complete. Either both component parts may be obtained in a separate
state, or else only one is separated and the other is converted into a
new compound in which its amount may be determined by weighing. This
will be a reaction of substitution, such as is often taken advantage of
for analysis. The first analysis of water was thus conducted in 1784
by Lavoisier and Meusnier. The apparatus they arranged consisted of a
glass retort containing water previously purified, and of which the
weight had been determined. The neck of the retort was inserted into
a porcelain tube, placed inside an oven, and heated to a red heat by
charcoal. Iron filings, which decompose water at a red heat, were placed
inside this tube. The end of the tube was connected with a worm, for
condensing any water which might pass through the tube undecomposed.
This condensed water was collected in a separate flask. The gas formed
by the decomposition was collected over water in a bell jar. The aqueous
vapour in passing over the red-hot iron was decomposed, and a gas was
formed from it whose weight could be determined from its volume, its
density being known. Besides the water which passed through the tube
unaltered, a certain quantity of water disappeared in the experiment,
and this quantity, in the experiments of Lavoisier and Meusnier, was
equal to the weight of gas which was collected in the bell jar plus the
increase in weight of the iron filings. Hence the water was decomposed
into a gas, which was collected in the bell jar, and a substance, which
combined with the iron; consequently, it is composed of these two
component parts. This was the first analysis of water ever made; but
here only one (and not both) of the gaseous component parts of water was
collected separately. Both the component parts of water can, however,
be simultaneously obtained in a free state. For this purpose the
decomposition is brought about by a galvanic current or by heat, as we
shall learn directly.[1]
[1] The first experiments of the synthesis and decomposition of water
did not afford, however, an entirely convincing proof that water
was composed of hydrogen and oxygen only. Davy, who investigated
the decomposition of water by the galvanic current, thought for a
long time that, besides the gases, an acid and alkali were also
obtained. He was only convinced of the fact that water contains
nothing but hydrogen and oxygen by a long series of researches,
which showed him that the appearance of an acid and alkali in the
decomposition of water proceeds from the presence of impurities
(especially from the presence of ammonium nitrate) in water. A
final comprehension of the composition of water is obtained from
the accurate determination of the quantities of the component
parts which enter into its composition. It will be seen from
this how many data are necessary for proving the composition of
water--that is, of the transformations of which it is capable.
What has been said of water refers to all other compounds; the
investigation of each one, the entire proof of its composition,
can only be obtained by the accumulation of a large mass of data
referring to it.
Water is a bad conductor of electricity--that is, pure water does not
transmit a feeble current; but if any salt or acid be dissolved in it,
then its conductivity increases, and _on the passage of a current_
through acidified water _it is decomposed_ into its component parts.
Some sulphuric acid is generally added to the water. By immersing
platinum plates (electrodes) in this water (platinum is chosen because
it is not acted on by acids, whilst many other metals are chemically
acted on by acids), and connecting them with a galvanic battery, it
will be observed that bubbles of gas appear on these plates. The gas
which separates is called _detonating gas_,[2] because, on ignition,
it very easily explodes.[3] What takes place is as follows:--First,
the water, by the action of the current, is decomposed into two gases.
The mixture of these gases forms detonating gas. When detonating gas
is brought into contact with an incandescent substance--for instance,
a lighted taper--the gases re-combine, forming water, the combination
being accompanied by a great evolution of heat, and therefore the vapour
of the water formed expands considerably, which it does very rapidly,
and as a consequence, an explosion takes place--that is, sound and
increase of pressure, and atmospheric disturbance, as in the explosion
of gunpowder.
[2] This gas is collected in a voltameter.
[3] In order to observe this explosion without the slightest danger,
it is best to proceed in the following manner. Some soapy water is
prepared, so that it easily forms soap bubbles, and it is poured
into an iron trough. In this water, the end of a gas-conducting
tube is immersed. This tube is connected with any suitable
apparatus, in which detonating gas is evolved. Soap bubbles,
full of this gas, are then formed. If the apparatus in which the
gas is produced be then removed (otherwise the explosion might
travel into the interior of the apparatus), and a lighted taper be
brought to the soap bubbles, a very sharp explosion takes place.
The bubbles should be small to avoid any danger; ten, each about
the size of a pea, suffice to give a sharp report, like a pistol
shot.
In order to discover what gases are obtained by the decomposition of
water, the gases which separate at each electrode must be collected
separately. For this purpose a V-shaped tube is taken; one of its ends
is open and the other fused up. A platinum wire, terminating inside
the tube in a plate, is fused into the closed end; the closed end is
entirely filled with water[4] acidified with sulphuric acid, and another
platinum wire, terminating in a plate, is immersed in the open end. If
a current from a galvanic battery be now passed through the wires an
evolution of gases will be observed, and the gas which is obtained in
the open branch passes into the air, while that in the closed branch
accumulates above the water. As this gas accumulates it displaces the
water, which continues to descend in the closed and ascend into the open
branch of the tubes. When the water, in this way, reaches the top of the
open end, the passage of the current is stopped, and the gas which was
evolved from one of the electrodes only is obtained in the apparatus.
By this means it is easy to prove that a particular gas appears at each
electrode. If the closed end be connected with the negative pole--_i.e._
with that joined to the zinc--then the gas collected in the apparatus
is capable of burning. This may be demonstrated by the following
experiment:--The bent tube is taken off the stand, and its open end
stopped up with the thumb and inclined in such a manner that the gas
passes from the closed to the open end. It will then be found, on
applying a lighted lamp or taper, that the gas burns. This combustible
gas is _hydrogen_. If the same experiment be carried on with a current
passing in the opposite direction--that is, if the closed end be joined
up with the positive pole (_i.e._ with the carbon, copper, or platinum),
then the gas which is evolved from it does not itself burn, but it
supports combustion very vigorously, so that a smouldering taper in it
immediately bursts into flame. This gas, which is collected at the anode
or positive pole, is _oxygen_, which is obtained, as we saw before (in
the Introduction), from mercury oxide and is contained in air.
[4] In order to fill the tube with water, it is turned up, so that
the closed end points downwards and the open end upwards, and
water acidified with sulphuric acid is poured into it.
Thus in the decomposition of water oxygen appears at the positive pole
and hydrogen at the negative pole,[4 bis] so that detonating gas will be
a mixture of both. Hydrogen burns in air from the fact that in doing so
it re-forms water, with the oxygen of the air. Detonating gas explodes
from the fact that the hydrogen burns in the oxygen mixed with it. It
is very easy to measure the relative quantities of one and the other
gas which are evolved in the decomposition of water. For this purpose
a funnel is taken, whose orifice is closed by a cork through which two
platinum wires pass. These wires are connected with a battery. Acidified
water is poured into the funnel, and a glass cylinder full of water
is placed over the end of each wire (fig. 18). On passing a current,
hydrogen and oxygen collect in these cylinders, and it will easily be
seen that two volumes of hydrogen are evolved for every one volume of
oxygen. This signifies that, in decomposing, water gives two volumes of
hydrogen and one volume of oxygen.
[4 bis] Owing to the gradual but steady progress made during the
last twenty-five years in the production of an electric current
from the dynamo and its transmission over considerable distances,
the electrolytic decomposition of many compound bodies has
acquired great importance, and the use of the electric current is
making its way into many chemical manufactures. Hence, Prof. D.
A. Lachinoff's proposal to obtain hydrogen and oxygen (both of
which have many applications) by means of electrolysis (either of
a 10 to 15 per cent. solution of caustic soda or a 15 per cent.
solution of sulphuric acid) may find a practical application, at
all events in the future. In general, owing to their simplicity,
electrolytic methods have a great future, but as yet, so long
as the production of an electric current remains so costly,
their application is limited. And for this reason, although
certain of these methods are mentioned in this work, they are
not specially considered, the more so since a profitable and
proper use of the electric current for chemical purposes requires
special electro-technical knowledge which beginners cannot he
assumed to have, and therefore, an exposition of the principles
of electrotechnology as applied, to the production of chemical
transformations, although referred to in places, does not come
within the scope of the present work.
[Illustration: FIG. 18.--Decomposition of water by the galvanic current,
for determining the relation between the volumes of hydrogen and oxygen.]
Water is also decomposed into its component parts by _the action of
heat_. At the melting point of silver (960°), and in its presence,
water is decomposed and the oxygen absorbed by the molten silver,
which dissolves it so long as it is liquid. But directly the silver
solidifies the oxygen is expelled from it. However, this experiment
is not entirely convincing; it might be thought that in this case the
decomposition of the water did not proceed from the action of heat, but
from the action of the silver on water--that silver decomposes water,
taking up the oxygen. If steam be passed through a red-hot tube, whose
internal temperature attains 1,000°, then a portion[5] of the water
decomposes into its component parts, forming detonating gas. But on
passing into the cooler portions of the apparatus this detonating gas
again reunites and forms water. The hydrogen and oxygen obtained combine
together at a lower temperature.[6] Apparently the problem--to show the
decomposability of water at high temperatures--is unattainable. It was
considered as such before Henri Sainte-Claire Deville (in the fifties)
introduced the conception of dissociation into chemistry, as of a change
of chemical state resembling evaporation, if decomposition be likened to
boiling, and before he had demonstrated the decomposability of water by
the action of heat in an experiment which will presently be described.
In order to demonstrate clearly the _dissociation_ of water, or its
decomposability by heat, at a temperature approaching that at which it
is formed, it was necessary to separate the hydrogen from the oxygen at
a high temperature, without allowing the mixture to cool. Deville took
advantage of the difference between the densities of hydrogen and oxygen.
[5] As water is formed by the combination of oxygen and hydrogen, with
a considerable evolution of heat, and as it can also be
decomposed, this reaction is a reversible one (_see_
Introduction), and consequently at a high temperature the
decomposition of water cannot be complete--it is limited by the
opposite reaction. Strictly speaking, it is not known how much
water is decomposed at a given temperature, although many efforts
(Bunsen, and others) have been made in various directions to solve
this question. Not knowing the coefficient of expansion, and the
specific heat of gases at such high temperatures, renders all
calculations (from observations of the pressure on explosion)
doubtful.
[6] Grove, in 1847, observed that a platinum wire fused in the
oxyhydrogen flame--that is, having acquired the temperature
of the formation of water--and having formed a molten drop at
its end which fell into water, evolved detonating gas--that
is, decomposed water. It therefore follows that water already
decomposes at the temperature of its formation. At that time, this
formed a scientific paradox; this we shall unravel only with the
development of the conceptions of dissociation, introduced into
science by Henri Sainte-Claire Deville, in 1857. These conceptions
form an important epoch in science, and their development is one
of the problems of modern chemistry. The essence of the matter is
that, at high temperatures, water exists but also decomposes, just
as a volatile liquid, at a certain temperature, exists both as a
liquid and as a vapour. Similarly as a volatile liquid saturates
a space, attaining its maximum tension, so also the products of
dissociation have their maximum tension, and once that is attained
decomposition ceases, just as evaporation ceases. Under like
conditions, if the vapour be allowed to escape (and therefore its
partial pressure be diminished), evaporation recommences, so also
if the products of decomposition be removed, decomposition again
continues. These simple conceptions of dissociation introduce
infinitely varied consequences into the mechanism of chemical
reactions, and therefore we shall have occasion to return to
them very often. We may add that Grove also concluded that water
was decomposed at a white heat, from the fact that he obtained
detonating gas by passing steam through a tube with a wire heated
strongly by an electric current, and also by passing steam over
molten oxide of lead, he obtained, on the one hand, litharge (=
oxide of lead and oxygen), and on the other, metallic lead formed
by the action of hydrogen.
[Illustration: FIG. 19.--Decomposition of water by the action of heat,
and the separation of the hydrogen formed by its permeating through a
porous tube.]
A wide porcelain tube P (fig. 19) is placed in a furnace, which can be
raised to a high temperature (it should be heated with small pieces of
good coke). In this tube there is inserted a second tube T, of smaller
diameter, made of unglazed earthenware and therefore porous. The ends
of the tube are luted to the wide tube, and two tubes, C and C', are
inserted into the ends, as shown in the drawing. With this arrangement
it is possible for a gas to pass into the annular space between the
walls of the two tubes, from whence it can be collected. Steam from a
retort or flask is passed through the tube D, into the inner porous
tube T. This steam on entering the red-hot space is decomposed into
hydrogen and oxygen. The densities of these gases are very different,
hydrogen being sixteen times lighter than oxygen. Light gases, as we
saw above, penetrate through porous surfaces very much more rapidly
than denser gases, and therefore the hydrogen passes through the
pores of the tube into the annular space very much more rapidly than
the oxygen. The hydrogen which separates out into the annular space
can only be collected when this space does not contain any oxygen.
If any air remains in this space, then the hydrogen which separates
out will combine with its oxygen and form water. For this reason a
gas incapable of supporting combustion--for instance, nitrogen or
carbonic anhydride--is previously passed into the annular space. Thus
the carbonic anhydride is passed through the tube C, and the hydrogen,
separated from the steam, is collected through the tube C', and will be
partly mixed with carbonic anhydride. A certain portion of the carbonic
anhydride will penetrate through the pores of the unglazed tube into
the interior of the tube T. The oxygen will remain in this tube, and
the volume of the remaining oxygen will be half that of the volume of
hydrogen which separates out from the annular space.[6 bis]
[6 bis] Part of the oxygen will also penetrate through the pores
of the tube; but, as was said before, a much smaller quantity
than the hydrogen, and as the density of oxygen is sixteen times
greater than that of hydrogen, the volume of oxygen which passes
through the porous walls will be four times less than the volume
of hydrogen (the quantities of gases passing through porous
walls are inversely proportional to the square roots of their
densities). The oxygen which separates out into the annular
space will combine, at a certain fall of temperature, with the
hydrogen; but as each volume of oxygen only requires two volumes
of hydrogen, whilst at least four volumes of hydrogen will pass
through the porous walls for every volume of oxygen that passes,
therefore, part of the hydrogen will remain free, and can be
collected from the annular space. A corresponding quantity of
oxygen remaining from the decomposition of the water can be
collected from the internal tube.
The decomposition of water is effected much more easily by a method of
substitution, taking advantage of the affinity of substances for the
oxygen or the hydrogen of water. If a substance be added to water,
which takes up the oxygen and replaces the hydrogen--then we shall
obtain the latter gas from the water. Thus with sodium, water gives
hydrogen, and with chlorine, which takes up the hydrogen, oxygen is
obtained.
Hydrogen is evolved from water by many metals, which are capable
of forming oxides in air--that is, which are capable of burning or
combining with oxygen. The capacity of metals for combining with
oxygen, and therefore for decomposing water, or for the evolution of
hydrogen, is very dissimilar.[7] Among metals, potassium and sodium
exhibit considerable energy in this respect. The first occurs in
potash, the second in soda. They are both lighter than water, soft, and
easily change in air. By bringing one or the other of them in contact
with water at the ordinary temperature,[8] a quantity of hydrogen,
corresponding with the amount of the metal taken, may be directly
obtained. One gram of hydrogen, occupying a volume of 11·16 litres at 0°
and 760 mm., is evolved from every 39 grams of potassium, or 23 grams of
sodium. The phenomenon may be observed in the following way: a solution
of sodium in mercury--or 'sodium amalgam,' as it is generally called--is
poured into a vessel containing water, and owing to its weight sinks to
the bottom; the sodium held in the mercury then acts on the water like
pure sodium, liberating hydrogen. The mercury does not act here, and the
same amount of it as was taken for dissolving the sodium is obtained in
the residue. The hydrogen is evolved gradually in the form of bubbles,
which pass through the liquid.
[7] In order to demonstrate the difference of the affinity of oxygen
for different elements, it is enough to compare the amounts of
heats which are evolved in their combination with 16 parts by
weight of oxygen; in the case of sodium (when Na_{2}O is formed,
or 46 parts of Na combine with 16 parts of oxygen, according to
Beketoff) 100,000 calories (or units of heat), are evolved, for
hydrogen (when water, H_{2}O, is formed) 69,000 calories, for
iron (when the oxide FeO is formed) 69,000, and if the oxide
Fe_{2}O_{3} is formed, 64,000 calories, for zinc (ZnO is formed)
86,000 calories, for lead (when PbO is formed) 51,000 calories,
for copper (when CuO is formed) 38,000 calories, and for mercury
(HgO is formed) 31,000 calories.
These figures cannot correspond directly with the magnitude of the
affinities, for the physical and mechanical side of the matter is
very different in the different cases. Hydrogen is a gas, and, in
combining with oxygen, gives a liquid; consequently it changes
its physical state, and, in doing so, evolves heat. But zinc and
copper are solids, and, in combining with oxygen, give solid
oxides. The oxygen, previously a gas, now passes into a solid or
liquid state, and, therefore, also must have given up its store of
heat in forming oxides. As we shall afterwards see, the degree of
contraction (and consequently of mechanical work) was different
in the different cases, and therefore the figures expressing the
heat of combination cannot directly depend on the affinities,
on the loss of internal energy previously in the elements.
Nevertheless, the figures above cited correspond, in a certain
degree, with the order in which the elements stand in respect
to their affinity for oxygen, as may be seen from the fact that
the mercury oxide, which evolves the least heat (among the above
examples), is the least stable is easily decomposed, giving up its
oxygen; whilst sodium, the formation of whose oxide is accompanied
by the greatest evolution of heat, is able to decompose all the
other oxides, taking up their oxygen. In order to generalise the
connection between affinity and the evolution and the absorption
of heat, which is evident in its general features, and was firmly
established by the researches of Favre and Silbermann (about
1840), and then of Thomsen (in Denmark) and Berthelot (in France),
many investigators, especially the one last mentioned, established
the _law of maximum work_. This states that only those chemical
reactions take place of their own accord in which the greatest
amount of chemical (latent, potential) energy is transformed into
heat. But, in the first place, we are not able, judging from what
has been said above, to distinguish that heat which corresponds
with purely chemical action from the sum total of the heat
observed in a reaction (in the calorimeter); in the second place,
there are evidently endothermal reactions which proceed under
the same circumstances as exothermal (carbon burns in the vapour
of sulphur with absorption of heat, whilst in oxygen it evolves
heat); and, in the third place, there are reversible reactions,
which when taking place in one direction evolve heat, and when
taking place in the opposite direction absorb it; and, therefore,
the principle of maximum work in its elementary form is not
supported by science. But the subject continues to be developed,
and will probably lead to a general law, such as thermal chemistry
does not at present possess.
[8] If a piece of metallic sodium be thrown into water, it floats on it
(owing to its lightness), keeps in a state of continual motion
(owing to the evolution of hydrogen on all sides), and immediately
decomposes the water, evolving hydrogen, which can be lighted.
This experiment may, however, lead to an explosion should the
sodium stick to the walls of the vessel, and begin to act on the
limited mass of water immediately adjacent to it (probably in
this case NaHO forms with Na, Na_{2}O, which acts on the water,
evolving much heat and rapidly forming steam), and the experiment
should therefore be carried on with caution. The decomposition
of water by sodium may he better demonstrated, and with greater
safety, in the following manner. Into a glass cylinder filled
with mercury, and immersed in a mercury bath, water is first
introduced, which will, owing to its lightness, rise to the top,
and then a piece of sodium wrapped in paper is introduced with
forceps into the cylinder. The metal rises through the mercury
to the surface of the water, on which it remains, and evolves
hydrogen, which collects in the cylinder, and may be tested after
the experiment has been completed. The safest method of making
this experiment is, however, as follows. The sodium (cleaned from
the naphtha in which it is kept) is either wrapped in fine copper
gauze and held by forceps, or else held in forceps at the end of
which a small copper cage is attached, and is then held under
water. The evolution of hydrogen goes on quietly, and it may he
collected in a bell jar and then lighted.
Beyond the hydrogen evolved and a solid substance, which remains in
solution (it may be obtained by evaporating the resultant solution) no
other products are here obtained. Consequently, from the two substances
(water and sodium) taken, the same number of new substances (hydrogen
and the substance dissolved in water) have been obtained, from which we
may conclude that the reaction which here takes place is a reaction of
double decomposition or of substitution. The resultant solid is nothing
else but the so-called caustic soda (sodium hydroxide), which is made
up of sodium, oxygen, and half of the hydrogen contained in the water.
Therefore, the substitution took place between the hydrogen and the
sodium, namely half of the hydrogen in the water was replaced by the
sodium, and was evolved in a free state. Hence the reaction which takes
place here may be expressed by the equation H_{2}O + Na = NaHO + H; the
meaning of this is clear from what has already been said.[9]
[9] This reaction is vigorously exothermal, _i.e._ it is accompanied
by the evolution of heat. If a sufficient quantity of water
be taken the whole of the sodium hydroxide, NaHO, formed is
dissolved, and about 42,500 units of heat are evolved per 23 grams
of sodium taken. As 40 grams of sodium hydroxide are produced,
and they in dissolving, judging from direct experiment, evolve
about 10,000 calories; therefore, without an excess of water, and
without the formation of a solution, the reaction would evolve
about 32,500 calories. We shall afterwards learn that hydrogen
contains in its smallest isolable particles H_{2} and not H,
and therefore it follows that the reaction should be written
thus--2Na + 2H_{2}O = H_{2} + 2NaOH, and it then corresponds with
an evolution of heat of +65,000 calories. And as N. N. Beketoff
showed that Na_{2}O, or anhydrous oxide of sodium, forms the
hydrate, or sodium hydroxide (caustic soda), 2NaHO, with water,
evolving about 35,500 calories, therefore the reaction 2Na +
H_{2}O = H_{2} + Na_{2}O corresponds to 29,500 calories. This
quantity of heat is less than that which is evolved in combining
with water, in the formation of caustic soda, and therefore it is
not to be wondered at that the hydrate, NaHO, is always formed
and not the anhydrous substance Na_{2}O. That such a conclusion,
which agrees with facts, is inevitable is also seen from the fact
that, according to Beketoff, the anhydrous sodium oxide, Na_{2}O,
acts directly on hydrogen, with separation of sodium, Na_{2}O + H
= NaHO + Na. This reaction is accompanied by an evolution of heat
equal to about 3,000 calories, because Na_{2}O + H_{2}O gives, as
we saw, 35,500 calories and Na + H_{2}O evolves 32,500 calories.
However, an opposite reaction also takes place--NaHO + Na =
Na_{2}O + H (both with the aid of heat)--consequently, in this
case heat is absorbed. In this we see an example of calorimetric
calculations and the limited application of the law of maximum
work for the general phenomena of reversible reactions, to which
the case just considered belongs. But it must be remarked that
all reversible reactions evolve or absorb but little heat, and
the reason of the law of maximum work, not being universal must
first of all be looked for in the fact that we have no means of
separating the heat which corresponds with the purely chemical
process from the sum total of the heat observed, and as the
structure of a number of substances is altered by heat and also
by contact, we can scarcely hope that the time approaches when
such a distinction will be possible. A heated substance, in point
of fact, has no longer the original energy of its atoms--that is,
the act of heating not only alters the store of motion of the
molecules but also of the atoms forming the molecules, in other
words, it makes the beginning of or preparation for chemical
change. From this it must be concluded that thermochemistry, or
the study of the heat accompanying chemical transformations,
cannot he identified with chemical mechanics. Thermo-chemical data
form a part of it, but they alone cannot give it.
Sodium and potassium act on water at the ordinary temperature. Other
heavier metals only act on it with a rise of temperature, and then not
so rapidly or vigorously. Thus magnesium and calcium only liberate
hydrogen from water at its boiling point, and zinc and iron only a red
heat, whilst a whole series of heavy metals, such as copper, lead,
mercury, silver, gold, and platinum, do not in the least decompose water
at any temperature, and do not replace its hydrogen.
From this it is clear that hydrogen may be obtained by the decomposition
of steam by the action of iron (or zinc) with a rise of temperature.
The experiment is conducted in the following manner: pieces of iron
(filings, nails, &c.), are placed in a porcelain tube, which is then
subjected to a strong heat and steam passed through it. The steam,
coming into contact with the iron, gives up its oxygen to it, and thus
the hydrogen is set free and passes out at the other end of the tube
together with undecomposed steam. This method, which is historically
very significant,[10] is practically inconvenient, as it requires a
rather high temperature. Further, this reaction, as a reversible one
(a red-hot mass of iron decomposes a current of steam, forming oxide
and hydrogen; and a mass of oxide of iron, heated to redness in a
stream of hydrogen, forms iron and steam), does not proceed in virtue
of the comparatively small difference between the affinity of oxygen
for iron (or zinc) and for hydrogen, but only because the hydrogen
escapes, as it is formed, in virtue of its elasticity.[11] If the
oxygen compounds--that is, the oxides--which are obtained from the
iron or zinc, be able to pass into solution, then the affinity acting
in solution is added, and the reaction may become non-reversible, and
proceed with comparatively much greater facility.[12] As the oxides
of iron and zinc, by themselves insoluble in water, are capable
of combining with (have an affinity for) acid oxides (as we shall
afterwards fully consider), and form saline and soluble substances, with
acids, or hydrates having acid properties, hence by the action of such
hydrates, or of their aqueous solutions,[13] iron and zinc are able to
liberate hydrogen with great ease at the ordinary temperature--that
is, they act on solutions of acids just as sodium acts on water.[14]
Sulphuric acid, H_{2}SO_{4}, is usually chosen for this purpose; the
hydrogen is displaced from it by many metals with much greater facility
than directly from water, and such a displacement is accompanied by
the evolution of a large amount of heat.[15] When the hydrogen in
sulphuric acid is replaced by a metal, a substance is obtained which is
called a salt of sulphuric acid or a sulphate. Thus, by the action of
zinc on sulphuric acid, hydrogen and zinc sulphate ZnSO_{4},[15 bis]
are obtained. The latter is a solid substance, soluble in water. In
order that the action of the metal on the acid should go on regularly,
and to the end, it is necessary that the acid should be diluted with
water, which dissolves the salt as it is formed; otherwise the salt
covers the metal, and hinders the acid from attacking it. Usually the
acid is diluted with from three to five times its volume of water, and
the metal is covered with this solution. In order that the metal should
act rapidly on the acid, it should present a large surface, so that a
maximum amount of the reacting substances may come into contact in a
given time. For this purpose the zinc is used as strips of sheet zinc,
or in the granulated form (that is, zinc which has been poured from a
certain height, in a molten state, into water). The iron should be in
the form of wire, nails, filings, or cuttings.
[Illustration: FIG. 20.--Apparatus for the preparation of hydrogen from
zinc and sulphuric acid.]
[10] The composition of water, as we saw above, was determined by
passing steam over red-hot iron; the same method has been used
for making hydrogen for filling balloons. An oxide having the
composition Fe_{3}O_{4} is formed in the reaction, so that it is
expressed by the equation 3Fe + 4H_{2}O = Fe_{3}O_{4} + 8H.
[11] The reaction between iron and water (note 10) is reversible. By
heating the oxide in a current of hydrogen, water and iron are
obtained. From this it follows, from the principle of chemical
equilibria, that if iron and hydrogen be taken, and also oxygen,
but in such a quantity that it is insufficient for combination
with both substances, then it will divide itself between the two;
part of it will combine with the iron and the other part with
the hydrogen, but a portion of both will remain in an uncombined
state.
Therefore, if iron and water be placed in a closed space,
decomposition of the water will proceed on heating to the
temperature at which the reaction 3Fe + 4H_{2}O = Fe_{3}O_{4} +
8H commences; but it ceases, does not go on to the end, because
the conditions for a reverse reaction are attained, and a state
of equilibrium will ensue after the decomposition of a certain
quantity of water. Here again (_see_ note 9) the reversibility is
connected with the small heat effect, and again both reactions
(direct and reverse) proceed at a red heat. But if, in the
above-described reaction, the hydrogen escapes as it is evolved,
then its partial pressure does not increase with its formation,
and therefore all the iron can he oxidised by the water. In
this we see the elements of that influence of mass to which we
shall have occasion to return later. With copper and lead there
will be no decomposition, either at the ordinary or at a high
temperature, because the affinity of these metals for oxygen is
much less than that of hydrogen.
[12] In general, if reversible as well as non-reversible reactions
can take place between substances acting on each other, then,
judging by our present knowledge, the non-reversible reactions
take place in the majority of cases, which obliges one to
acknowledge the action, in this case, of comparatively strong
affinities. The reaction, Zn + H_{2}SO_{4} = H_{2} + ZnSO_{4},
which takes place in solutions at the ordinary temperature, is
scarcely reversible under these conditions, but at a certain high
temperature it becomes reversible, because at this temperature
zinc sulphate and sulphuric acid split up, and the action must
take place between the water and zinc. From the preceding
proposition results proceed which are in some cases verified
by experiment. If the action of zinc or iron on a solution of
sulphuric acid presents a non-reversible reaction, then we may
by this means obtain hydrogen in a very compressed state, and
compressed hydrogen will not act on solutions of sulphates of
the above-named metals. This is verified in reality as far as
was possible in the experiments to keep up the compression or
pressure of the hydrogen. Those metals which do not evolve
hydrogen with acids, on the contrary, should, at least at an
increase of pressure, be displaced by hydrogen. And in fact
Brunner showed that gaseous hydrogen displaces platinum and
palladium from the aqueous solutions of their chlorine compounds,
but not gold, and Beketoff succeeded in showing that silver and
mercury, under a considerable pressure, are separated from the
solutions of certain of their compounds by means of hydrogen.
Reaction already commences under a pressure of six atmospheres,
if a weak solution of silver sulphate be taken; with a stronger
solution a much greater pressure is required, however, for the
separation of the silver.
[13] For the same reason, many metals in acting on solutions of the
alkalis displace hydrogen. Aluminium acts particularly clearly
in this respect, because its oxide gives a soluble compound with
alkalis. For the same reason tin, in acting on hydrochloric acid,
evolves hydrogen, and silicon does the same with hydrofluoric
acid. It is evident that in such cases the sum of all the
affinities plays a part; for instance, taking the action of
zinc on sulphuric acid, we have the affinity of zinc for oxygen
(forming zinc oxide, ZnO), the affinity of its oxide for
sulphuric anhydride, SO_{3} (forming zinc sulphate, ZnSO_{4}),
and the affinity of the resultant salt, ZnSO_{4}, for water.
It is only the first-named affinity that acts in the reaction
between water and the metal, if no account is taken of those
forces (of a physico-mechanical character) which act between the
molecules (for instance, the cohesion between the molecules of
the oxide) and those forces (of a chemical character) which act
between the atoms forming the molecule, for instance, between
the atoms of hydrogen giving the molecule H_{2} containing two
atoms. I consider it necessary to remark, that the hypothesis of
the affinity or endeavour of heterogeneous atoms to enter into a
common system and in harmonious motion (_i.e._ to form a compound
molecule) must inevitably be in accordance with the hypothesis of
forces including homogeneous atoms to form complex molecules (for
instance, H_{2}), and to build up the latter into solid or liquid
substances, in which the existence of an attraction between the
homogeneous particles must certainly be admitted. Therefore,
those forces which bring about solution must also be taken into
consideration. These are all forces of one and the same series,
and in this may be seen the great difficulties surrounding the
study of molecular mechanics and its province--chemical mechanics.
[14] It is acknowledged that zinc itself acts on water, even at the
ordinary temperature, but that the action is confined to small
masses and only proceeds at the surface. In reality, zinc, in the
form of a very fine powder, or so-called 'zinc dust,' is capable
of decomposing water with the formation of oxide (hydrated) and
hydrogen. The oxide formed acts on sulphuric acid, water then
dissolves the salt produced, and the action continues because
one of the products of the action of water on zinc, zinc oxide,
is removed from the surface. One might naturally imagine that
the reaction does not proceed directly between the metal and
water, but between the metal and the acid, but such a simple
representation, which we shall cite afterwards, hides the
mechanism of the reaction, and does not permit of its actual
complexity being seen.
[15] According to Thomsen the reaction between zinc and a very weak
solution of sulphuric acid evolves about 38,000 calories (zinc
sulphate being formed) per 65 parts by weight of zinc; and 56
parts by weight of iron--which combine, like 65 parts by weight
of zinc, with 16 parts by weight of oxygen--evolve about 25,000
calories (forming ferrous sulphate, FeSO_{4}). Paracelsus
observed the action of metals on acids in the seventeenth
century; but it was not until the eighteenth century that Lémery
determined that the gas which is evolved in this action is a
particular one which differs from air and is capable of burning.
Even Boyle confused it with air. Cavendish determined the chief
properties of the gas discovered by Paracelsus. At first it was
called 'inflammable air'; later, when it was recognised that in
burning it gives water, it was called hydrogen, from the Greek
words for water and generator.
[15 bis] If, when the sulphuric acid is poured over the zinc, the
evolution of the hydrogen proceed too slowly, it may be greatly
accelerated by adding a small quantity of a solution of CuSO_{4}
or PtCl_{4} to the acid. The reason of this is explained in Chap.
XVI., note 10 bis.
The usual method of obtaining hydrogen is as follows:--A certain
quantity of granulated zinc is put into a double-necked, or Woulfe's,
bottle. Into one neck a funnel is placed, reaching to the bottom of
the bottle, so that the liquid poured in may prevent the hydrogen from
escaping through it. The gas escapes through a special gas conducting
tube, which is firmly fixed, by a cork, into the other neck, and ends in
a water bath (fig. 20), under the orifice of a glass cylinder full of
water.[16] If sulphuric acid be now poured into the Woulfe's bottle it
will soon be seen that bubbles of a gas are evolved, which is hydrogen.
The first part of the gas evolved should not be collected, as it is
mixed with the air originally in the apparatus. This precaution should
be taken in the preparation of all gases. Time must be allowed for the
gas evolved to displace all the air from the apparatus, otherwise in
testing the combustibility of the hydrogen an explosion may occur from
the formation of detonating gas (the mixture of the oxygen of the air
with the hydrogen).[17]
[16] As laboratory experiments with gases require a certain preliminary
knowledge, we will describe certain _practical methods for the
collection and preparation of gases_. When in laboratory practice
an intermittent supply of hydrogen (or other gas which is evolved
without the aid of heat) is required the apparatus represented
in fig. 21 is the most convenient. It consists of two bottles,
having orifices at the bottom, in which corks with tubes are
placed, and these tubes are connected by an india-rubber tube
(sometimes furnished with a spring clamp). Zinc is placed in one
bottle, and dilute sulphuric acid in the other. The neck of the
former is closed by a cork, which is fitted with a gas-conducting
tube with a stopcock. If the two bottles are connected with each
other and the stopcock be opened, the acid will flow to the zinc
and evolve hydrogen. If the stopcock be closed, the hydrogen
will force out the acid from the bottle containing the zinc, and
the action will cease. Or the vessel containing the acid may be
placed at a lower level than that containing the zinc, when all
the liquid will flow into it, and in order to start the action
the acid vessel may be placed on a higher level than the other,
and the acid will flow to the zinc. It can also be employed for
collecting gases (as an aspirator or gasometer).
[Illustration: FIG. 21.--A very convenient apparatus for the
preparation of gases obtained without heat. It may also replace
an aspirator or gasometer.]
In laboratory practice, however, other forms of apparatus are
generally employed for exhausting, collecting, and holding gases.
We will here cite the most usual forms. An _aspirator_ usually
consists of a vessel furnished with a stopcock at the bottom.
A stout cork, through which a glass tube passes, is fixed into
the neck of this vessel. If the vessel be filled up with water
to the cork and the bottom stopcock is opened, then the water
will run out and draw gas in. For this purpose the glass tube is
connected with the apparatus from which it is desired to pump out
or exhaust the gas.
[Illustration: FIG. 22.--Continuous aspirator. The tube _d_
should be more than 32 feet long.]
The aspirator represented in fig. 22 may be recommended for its
continuous action. It consists of a tube _d_ which widens out at
the top, the lower part being long and narrow. In the expanded
upper portion _c_, two tubes are sealed; one, _e_, for drawing in
the gas, whilst the other, _b_, is connected to the water supply
_w_. The amount of water supplied through the tube _b_ must be
less than the amount which can be carried off by the tube _d_.
Owing to this the water in the tube _d_ will flow through it in
cylinders alternating with cylinders of gas, which will be thus
carried away. The gas which is drawn through may be collected
from the end of the tube _d_, but this form of pump is usually
employed where the air or gas aspirated is not to be collected.
If the tube _d_ is of considerable length, say 40 ft. or more,
a very fair vacuum will be produced, the amount of which is
shown by the gauge _g_; it is often used for filtering under
reduced pressure, as shown in the figure. If water be replaced by
mercury, and the length of the tube _d_ be greater than 760 mm.,
the aspirator may be employed as an air-pump, and all the air may
be exhausted from a limited space; for instance, by connecting
_g_ with a hollow sphere.
[Illustration: FIG. 23.--Gasholder.]
_Gasholders_ are often used for collecting and holding gases.
They are made of glass, copper, or tin plate. The usual form
is shown in fig. 23. The lower vessel _B_ is made hermetically
tight--_i.e._, impervious to gases--and is filled with water.
A funnel is attached to this vessel (on several supports).
The vessel _B_ communicates with the bottom of the funnel by
a stopcock _b_ and a tube _a_, reaching to the bottom of the
vessel _B_. If water be poured into the funnel and the stopcocks
_a_ and _b_ opened, the water will run through _a_, and the air
escape from the vessel _B_ by _b_. A glass tube _f_ runs up the
side of the vessel _B_, with which it communicates at the top
and bottom, and shows the amount of water and gas the gasholder
contains. In order to fill the gasholder with a gas, it is first
filled with water, the cocks _a_, _b_ and _e_ are closed, the nut
_d_ unscrewed, and the end of the tube conducting the gas from
the apparatus in which it is generated is passed into _d_. As
the gas fills the gasholder, the water runs out at _d_. If the
pressure of a gas be not greater than the atmospheric pressure
and it be required to collect it in the gasholder, then the
stopcock _e_ is put into communication with the space containing
the gas. Then, having opened the orifice _d_, the gasholder
acts like an aspirator; the gas will pass through _e_, and the
water run out at _d_. If the cocks be closed, the gas collected
in the gasholder may be easily preserved and transported. If
it be desired to transfer this gas into another vessel, then a
gas-conducting tube is attached to _e_, the cock _a_ opened, _b_
and _d_ closed, and the gas will then pass out at _e_, owing to
its pressure in the apparatus being greater than the atmospheric
pressure, due to the pressure of the water poured into the
funnel. If it be required to fill a cylinder or flask with the
gas, it is filled with water and inverted in the funnel, and the
stopcocks _b_ and _a_ opened. Then water will run through _a_,
and the gas will escape from the gasholder into the cylinder
through _b_.
[17] When it is required to prepare hydrogen in large quantities for
filling balloons, copper vessels or wooden casks lined with
lead are employed; they are filled with scrap iron, over which
dilute sulphuric acid is poured. The hydrogen generated from a
number of casks is carried through lead pipes into special casks
containing water (in order to cool the gas) and lime (in order
to remove acid fumes). To avoid loss of gas all the joints are
made hermetically tight with cement or tar. In order to fill his
gigantic balloon (of 25,000 cubic metres capacity), Giffard, in
1878, constructed a complicated apparatus for giving a continuous
supply of hydrogen, in which a mixture of sulphuric acid and
water was continually run into vessels containing iron, and
from which the solution of iron sulphate formed was continually
drawn off. When coal gas, extracted from coal, is employed for
filling balloons, it should be as light, or as rich in hydrogen,
as possible. For this reason, only the last portions of the
gas coming from the retorts are collected, and, besides this,
it is then sometimes passed through red-hot vessels, in order
to decompose the hydrocarbons as much as possible; charcoal
is deposited in the red-hot vessels, and hydrogen remains as
gas. Coal gas may be yet further enriched in hydrogen, and
consequently rendered lighter, by passing it over an ignited
mixture of charcoal and lime.
L. Mond (London) proposes to manufacture hydrogen on a large
scale from water gas (_see infra_, and Chapters VIII. and IX.),
which contains a mixture of oxide of carbon (CO) and hydrogen,
and is produced by the action of steam upon incandescent coke
(C + H_{2}O = CO + H_{2}). He destroys the oxide of carbon
by converting it into carbon and carbonic anhydride (2CO =
C + CO_{2}), which is easily done by means of incandescent,
finely-divided metallic nickel; the carbon then remains with
the nickel, from which it may be removed by burning it in air,
and the nickel can then be used over again (_see_ Chapter IX.,
Note 24 bis). The CO_{2} formed is removed from the hydrogen by
passing it through milk of lime. This process should apparently
give hydrogen on a large scale more economically than any of the
methods hitherto proposed.
Hydrogen, besides being contained in water, is also contained in many
other substances,[18] and may be obtained from them. As examples of
this, it may be mentioned (1) that a mixture of formate of sodium,
CHNaO_{2}, and caustic soda, NaHO, when heated to redness, forms sodium
carbonate, Na_{2}CO_{3}, and hydrogen, H_{2};[19] (2) that a number of
organic substances are decomposed at a red heat, forming hydrogen, among
other gases, and thus it is that hydrogen is contained in ordinary coal
gas.
[18] Of the metals, only a very few combine with hydrogen (for example,
sodium), and give substances which are easily decomposed. Of
the non-metals, the halogens (fluorine, chlorine, bromine,
and iodine) most easily form hydrogen compounds; of these the
hydrogen compound of chlorine, and still more that of fluorine,
is stable, whilst those of bromine and iodine are easily
decomposed, especially the latter. The other non-metals--for
instance, sulphur, carbon, and phosphorus--give hydrogen
compounds of different composition and properties, but they are
all less stable than water. The number of the carbon compounds of
hydrogen is enormous, but there are very few among them which are
not decomposed, with separation of the carbon and hydrogen, at a
red heat.
[19] The reaction expressed by the equation CNaHO_{2} + NaHO =
CNa_{2}O_{3} + H_{2} may be effected in a glass vessel, like
the decomposition of copper carbonate or mercury oxide (_see_
Introduction); it is non-reversible, and takes place without the
presence of water, and therefore Pictet (_see_ later) made use of
it to obtain hydrogen under great pressure.
Charcoal itself liberates hydrogen from steam at a high temperature;[20]
but the reaction which here takes place is distinguished by a certain
complexity, and will therefore be considered later.
[20] The reaction between charcoal and superheated steam is a double
one--that is, there may be formed either carbonic oxide, CO
(according to the equation H_{2}O + C = H_{2} + CO), or carbonic
anhydride CO_{2} (according to the equation 2H_{2}O + C = 2H_{2}
+ CO_{2}), and the resulting mixture is called _water-gas_; we
shall speak of it in Chapter IX.
_The properties of hydrogen._--Hydrogen presents us with an example
of a gas which at first sight does not differ from air. It is not
surprising, therefore, that Paracelsus, having discovered that an
aëriform substance is obtained by the action of metals on sulphuric
acid, did not determine exactly its difference from air. In fact,
hydrogen, like air, is colourless, and has no smell;[21] but a more
intimate acquaintance with its properties proves it to be entirely
different from air. The first sign which distinguishes hydrogen from
air is its combustibility. This property is so easily observed that
it is the one to which recourse is usually had in order to recognise
hydrogen, if it is evolved in a reaction, although there are many other
combustible gases. But before speaking of the combustibility and other
chemical properties of hydrogen, we will first describe the physical
properties of this gas, as we did in the case of water. It is easy to
show that it is one of the lightest gases.[22] If passed into the bottom
of a flask full of air, hydrogen will not remain in it, but, owing
to its lightness, rapidly escapes and mixes with the atmosphere. If,
however, a cylinder whose orifice is turned downwards be filled with
hydrogen, it will not escape, or, more correctly, it will only slowly
mix with the atmosphere. This may be demonstrated by the fact that a
lighted taper sets fire to the hydrogen at the orifice of the cylinder,
and is itself extinguished inside the cylinder. Hence, hydrogen, being
itself combustible, does not support combustion. The great lightness
of hydrogen is taken advantage of for balloons. Ordinary coal gas,
which is often also used for the same purpose, is only about twice
as light as air, whilst hydrogen is 14-1/2 times lighter than air.
A very simple experiment with soap bubbles very well illustrates the
application of hydrogen for filling balloons. Charles, of Paris, showed
the lightness of hydrogen in this way, and constructed a balloon filled
with hydrogen almost simultaneously with Montgolfier. One litre of pure
and dry hydrogen[23] at 0° and 760 mm. pressure weighs 0·08986 gram;
that is, hydrogen is almost 14-1/2 (more exactly, 14·39) times lighter
than air. It is the lightest of all gases. The small density of hydrogen
determines many remarkable properties which it shows; thus, hydrogen
passes exceedingly rapidly through fine orifices, its molecules (Chapter
I.) being endued with the greatest velocity.[24] At pressures somewhat
higher than the atmospheric pressure, all other gases exhibit a greater
compressibility and co-efficient of expansion than they should according
to the laws of Mariotte and Gay-Lussac; whilst hydrogen, on the
contrary, is compressed to a less degree than it should be from the law
of Mariotte,[25] and with a rise of pressure it expands slightly less
than at the atmospheric pressure.[26] However, hydrogen, like air and
many other gases which are permanent at the ordinary temperature, does
not pass into a liquid state under a very considerable pressure,[27] but
is compressed into a lesser volume than would follow from Mariotte's
law.[28] From this it may be concluded that the absolute boiling point
of hydrogen, and of gases resembling it,[29] lies very much below the
ordinary temperature; that is, that the liquefaction of this _gas_
is only possible at low temperatures, and under great pressures.[30]
This conclusion was verified (1877) by the experiments of Pictet
and Cailletet.[31] They compressed gases at a very low temperature,
and then allowed them to expand, either by directly decreasing the
pressure or by allowing them to escape into the air, by which means
the temperature fell still lower, and then, just as steam when rapidly
rarefied[32] deposits liquid water in the form of a fog, hydrogen in
expanding forms a fog, thus indicating its passage into a liquid state.
But as yet it has been impossible to preserve this liquid, even for a
short time, to determine its properties, notwithstanding the employment
of a temperature of -200° and a pressure of 200 atmospheres,[33] although
by these means the gases of the atmosphere may be kept in a liquid
state for a long time. This is due to the fact that the absolute boiling
point of hydrogen lies lower than that of all other known gases, which
also depends on the extreme lightness of hydrogen.[34]
[21] Hydrogen obtained by the action of zinc or iron on sulphuric acid
generally smells of hydrogen sulphide (like rotten eggs), which
it contains in admixture. As a rule such hydrogen is not so
pure as that obtained by the action of an electric current or
of sodium on water. The impurity of the hydrogen depends on the
impurities contained in the zinc, or iron, and sulphuric acid,
and on secondary reactions which take place simultaneously with
the main reaction. Impure hydrogen may be easily freed from
the impurities it contains: some of them--namely, those having
acid properties--are absorbed by caustic soda, and therefore
may be removed by passing the hydrogen through a solution of
this substance; another series of impurities is absorbed by a
solution of mercuric chloride; and, lastly, a third series is
absorbed by a solution of potassium permanganate. If absolutely
_pure hydrogen_ be required, it is sometimes obtained by the
decomposition of water (previously boiled to expel all air,
and mixed with pure sulphuric acid) by the galvanic current.
Only the gas evolved at the negative electrode is collected.
Or else, an apparatus like that which gives detonating gas is
used, the positive electrode, however, being immersed under
mercury containing zinc in solution. The oxygen which is evolved
at this electrode then immediately, at the moment of its
evolution, combines with the zinc, and this compound dissolves
in the sulphuric acid and forms zinc sulphate, which remains in
solution, and therefore the hydrogen generated will be quite free
from oxygen.
[22] An inverted beaker is attached to one arm of the beam of a
tolerably sensitive balance, and its weight counterpoised by
weights in the pan attached to the other arm, If the beaker
be then filled with hydrogen it rises, owing to the air being
replaced by hydrogen. Thus, at the ordinary temperature of a
room, a litre of air weighs about 1·2 gram, and on replacing the
air by hydrogen a decrease in weight of about 1 gram per litre is
obtained. Moist hydrogen is heavier than dry--for aqueous vapour
is nine times heavier than hydrogen. In filling balloons it is
usually calculated that (it being impossible to have perfectly
dry hydrogen or to obtain it quite free from air) the lifting
force due to the difference between the weights of equal volumes
of hydrogen and air is equal to 1 kilogram (= 1,000 grams) per
cubic metre (= 1,000 litres).
[23] The density of hydrogen in relation to the air has been repeatedly
determined by accurate experiments. The first determination, made
by Lavoisier, was not very exact; taking the density of air as
unity, he obtained 0·0769 for that of hydrogen--that is, hydrogen
as thirteen times lighter than air. More accurate determinations
are due to Thomsen, who obtained the figure 0·0693; Berzelius
and Dulong, who obtained 0·0688; and Dumas and Boussingault, who
obtained 0·06945. Regnault, and more recently Le Duc (1892),
took two spheres of considerable capacity, which contained equal
volumes of air (thus avoiding the necessity of any correction
for weighing them in air). Both spheres were attached to the
scale pans of a balance. One was sealed up, and the other first
weighed empty and then full of hydrogen. Thus, knowing the weight
of the hydrogen filling the sphere, and the capacity of the
sphere, it was easy to find the weight of a litre of hydrogen;
and, knowing the weight of a litre of air at the same temperature
and pressure, it was easy to calculate the density of hydrogen.
Regnault, by these experiments, found the average density of
hydrogen to be 0·06926 in relation to air; Le Duc, 0·06948 (with
a possible error of ±0·00001), and this latter figure must now be
looked upon as near to the truth.
In this work I shall always refer the densities of all gases to
hydrogen, and not to air; I will therefore give, for the sake of
clearness, the weight of a litre of dry pure hydrogen in grams
at a temperature _t_° and under a pressure _H_ (measured in
millimetres of mercury at 0°, in lat. 45°). The weight of a litre
of hydrogen
= 0·08986 × (_H_/760) × 1/(1 + 0·00367_t_) gram.
For aëronauts it is very useful to know, besides this, the
weight of the air at different heights, and I therefore insert
the adjoining table, constructed on the basis of Glaisher's
data, for the temperature and moisture of the atmospheric strata
in clear weather. All the figures are given in the metrical
system--1,000 millimetres = 39·37 inches, 1,000 kilograms =
2204·3375 lbs., 1,000 cubic metres = 35,316·6 cubic feet. The
starting temperature at the earth's surface is taken as = 15° C.,
its moisture 60 p.c., pressure 760 millimetres. The pressures
are taken as indicated by an _aneroid barometer_, assumed to be
corrected at the sea level and at lat. 45° C. If the height above
the level of the sea equal _z_ kilometres, then the weight of 1
cubic metre of air may be approximately taken as 1·222- 0·12_z_ +
0·00377_z_^2 kilogram.
+--------+-----------+--------+--------+--------------------+
|Pressure|Temperature|Moisture| Height | Weight of the air |
| | | |(metres)|(1,000 cubic metres)|
|--------+-----------+--------+--------+--------------------+
| 760 mm.| 15° C. | 60 p.c.| 0 | 1222 kilos. |
| 700 " | 11·0° " | 64 " | 690 | 1141 " |
| 650 " | 7·6° " | 64 " | 1300 | 1073 " |
| 600 " | 4·3° " | 63 " | 1960 | 1003 " |
| 550 " | -1·0° " | 62 " | 2660 | 931 " |
| 500 " | -2·4° " | 58 " | 3420 | 857 " |
| 450 " | -5·8° " | 52 " | 4250 | 781 " |
| 400 " | -9·1° " | 44 " | 5170 | 703 " |
| 350 " | -12·5° " | 36 " | 6190 | 624 " |
| 300 " | -15·9° " | 27 " | 7360 | 542 " |
| 250 " | -19·2° " | 18 " | 8720 | 457 " |
+--------+-----------+--------+--------+--------------------+
Although the figures in this table are calculated with every
possible care from average data, yet they can only be taken
approximately, for in every separate case the conditions, both at
the earth's surface and in the atmosphere, will differ from those
here taken. In calculating the height to which a balloon can
ascend, it is evident that the density of gas in relation to air
must be known. This density for ordinary coal gas is from 0·6 to
0·35, and for hydrogen with its ordinary contents of moisture and
air from 0·1 to 0·15.
Hence, for instance, it may be calculated that a balloon of 1,000
cubic metres capacity filled with pure hydrogen, and weighing
(the envelope, tackle, people, and ballast) 727 kilograms, will
only ascend to a height of about 4,250 metres.
[24] If a cracked flask be filled with hydrogen and its neck immersed
under water or mercury, then the liquid will rise up into the
flask, owing to the hydrogen passing through the cracks about 3·8
times quicker than the air is able to pass through these cracks
into the flask. The same phenomenon may be better observed if,
instead of a flask, a tube be employed, whose end is closed by a
porous substance, such as graphite, unglazed earthenware, or a
gypsum plate.
[25] According to Boyle and Mariotte's law, for a given gas at a
constant temperature the volume decreases by as many times as
the pressure increases; that is, this law requires that the
product of the volume _v_ and the pressure _p_ for a given gas
should be a constant quantity: _pv_ = _C_, a constant quantity
which does not vary with a change of pressure. This equation
does very nearly and exactly express the observed relation
between the volume and pressure, but only within comparatively
small variations of pressure, density, and volume. If these
variations be in any degree considerable, the quantity _pv_
proves to be dependent on the pressure, and it either increases
or diminishes with an increase of pressure. In the former case
the compressibility is less than it should he according to
Mariotte's law, in the latter case it is greater. We will call
the first case a positive discrepancy (because then _d(pv)/d(p)_
is greater than zero), and the second case a negative discrepancy
(because then _d(pv)/d(p)_ is less than zero). Determinations
made by myself (in the seventies), M. L. Kirpicheff, and V. A.
Hemilian showed that all known gases at low pressures--_i.e._
when considerably rarefied--present positive discrepancies. On
the other hand, it appears from the researches of Cailletet,
Natterer, and Amagat that all gases under great pressures (when
the volume obtained is 500-1,000 times less than under the
atmospheric pressure) also present positive discrepancies. Thus
under a pressure of 2,700 atmospheres air is compressed, not
2,700 times, but only 800, and hydrogen 1,000 times. Hence the
positive kind of discrepancy is, so to say, normal to gases. And
this is easily intelligible. If a gas followed Mariotte's law, or
if it were compressed to a greater extent than is shown by this
law, then under great pressures it would attain a density greater
than that of solid and liquid substances, which is in itself
improbable and even impossible by reason of the fact that solid
and liquid substances are themselves but little compressible.
For instance, a cubic centimetre of oxygen at 0° and under the
atmospheric pressure weighs about 0·0014 gram, and at a pressure
of 3,000 atmospheres (this pressure is attained in guns) it
would, if it followed Mariotte's law, weigh 4·2 grams--that is,
would be about four times heavier than water--and at a pressure
of 10,000 atmospheres it would be heavier than mercury. Besides
this, positive discrepancies are probable because the molecules
of a gas themselves must occupy a certain volume. Considering
that Mariotte's law, strictly speaking, applies only to the
intermolecular space, we can understand the necessity of positive
discrepancies. If we designate the volume of the molecules of a
gas by _b_ (like van der Waals, _see_ Chap. I., Note 34), then it
must be expected that _p(v-b) = C_. Hence _pv = C + bp_, which
expresses a positive discrepancy. Supposing that for hydrogen
_pv_ = 1,000, at a pressure of one metre of mercury, according to
the results of Regnault's, Amagat's, and Natterer's experiments,
we obtain _b_ as approximately 0·7 to 0·9.
Thus the increase of _pv_ with the increase of pressure must
be considered as the normal law of the compressibility of
gases. Hydrogen presents such a positive compressibility at
all pressures, for it presents positive discrepancies from
Mariotte's law, according to Regnault, at all pressures above the
atmospheric pressure. Hence hydrogen is, so to say, a perfect
gas. No other gas behaves so simply with a change of pressure.
All other gases at pressures from 1 to 30 atmospheres present
negative discrepancies--that is, they are then compressed to a
greater degree than should follow from Mariotte's law, as was
shown by the determinations of Regnault, which were verified
when repeated by myself and Boguzsky. Thus, for example, on
changing the pressure from 4 to 20 metres of mercury--that is,
on increasing the pressure five times--the volume only decreased
4·93 times when hydrogen was taken, and 5·06 when air was taken.