CHEMISTRY WITHOUT A LABORATORY.
We have already pointed out the possibility of going through a course
of physics without any special apparatus, we shall now endeavour to
show our readers the method of performing some experiments in chemistry
without a laboratory, or at any rate with only a few simple and
inexpensive appliances. The preparation of gases, such as hydrogen,
carbonic acid, and oxygen, is very easily accomplished, but we shall
here point out principally a series of experiments that are not so much
known. We will commence, for example, by describing an interesting
experiment which often occurs in a course of chemistry. Ammoniacal
gas combined with the elements of water is analogous to a metallic
oxide which includes a metallic root, _ammonium_. This hypothetically
composed metal may be in a manner perceived, since it is possible to
amalgamate it with mercury by operating in the following manner:—We
take a porcelain mortar, in which we pour a quantity of mercury, and
then cut some thin strips of sodium, which are thrown into the mercury.
By stirring it about with the pestle a loud cracking is produced,
accompanied by a flame, which bears evidence to the union of the
mercury and the sodium, and the formation of an amalgam of sodium. If
this amalgam of sodium is put into a slender glass tube containing a
concentrated solution of hydrochlorate of ammonia in water, we see
the ammonia expand in an extraordinary manner, and spout out from the
end of the tube, which is now too small to contain it, in the form
of a metallic substance (fig. 302). In this case, the ammonium, the
radical which exists in the ammoniacal salts, becomes amalgamated with
the mercury, driving out the sodium with which it had previously been
combined; the ammonium thus united with the mercury becomes decomposed
in ammoniacal gas and hydrogen, the mercury assuming its ordinary form.
Phosphate of ammonia is very valuable from its property of rendering
the lightest materials, such as gauze or muslin, incombustible. Dip a
piece of muslin in a solution of phosphate of ammonia, and dry it in
contact with the air; that done, you will find it is impossible to set
fire to the material; it will get charred, but you cannot make it burn.
It is to be wished that this useful precaution were oftener taken in
the matter of ball-dresses, which have so frequently been the cause of
serious accidents. There is no danger whatever with a dress that has
been soaked in phosphate of ammonia, which is very inexpensive, and
easily procured.
[Illustration: Fig. 302.—Experiment with ammonium.]
For preparing cool drinks in the summer ammoniacal salts are very
useful; some _nitrate of ammonia_ mixed with its weight in water,
produces a considerable lowering of the temperature, and is very useful
for making ice. _Volatile alkali_, which is so useful an application
for stings from insects, is a solution of ammoniacal gas in water, and
_sal-volatile_, which has such a refreshing and reanimating odour, is
a carbonate of ammonia. We often see in chemists’ shops large glass
jars, the insides of which are covered with beautiful transparent white
crystals, which are formed over a red powder placed at the bottom of
the vase. These crystals are the result of a combination of cyanogen
and iodine. Nothing is easier than the preparation of _iodide of
cyanogen_, a very volatile body, which possesses a strong tendency to
assume a definite crystalline form. We pound in a mortar a mixture of
50 grams of cyanide of mercury, and 100 grams of iodine; under the
action of the pestle the powder, which was at first a brownish colour,
assumes a shade of bright vermilion red. The cyanogen combines with
the iodine, and transforms itself into fumes with great rapidity. If
the powder is placed at the bottom of a stoppered glass jar, the fumes
of the iodide of cyanogen immediately condense, thereby producing
beautiful crystals which often attain considerable size (fig. 303).
Cyanogen forms with sulphur a remarkable substance, _sulpho-cyanogen_,
the properties of which we cannot describe without exceeding the
limits of our present treatise; we shall therefore confine ourselves
to pointing out one of its combinations, which is well known at the
present day, owing to its singular properties. This is sulpho-cyanide
of mercury, with which small combustible cones are made, generally
designated by the pompous title of _Pharaoh’s serpents_. For making
these, some sulpho-cyanide of potassium is poured into a solution of
nitrate acid on mercury, which forms a precipitate of sulpho-cyanide
of mercury. This is a white, combustible powder, which after passing
through a filter, should be transformed into a stiff pulp by means of
water containing a solution of gum. The pulp is afterwards mixed with
a small quantity of nitrate of potash, and fashioned into cones or
cylinders of about an inch and a quarter in length, which should be
thoroughly dried. The egg thus obtained can be hatched by the simple
application of a lighted match, and gives rise to the phenomenon.
The sulpho-cyanide slowly expands, the cylinder increases in length,
and changes to a yellowish substance, which dilates and extends to a
length of twenty or five-and-twenty inches. It has the appearance of a
genuine serpent, which has just started into existence, and stretches
out its tortuous coils, endeavouring to escape from its narrow prison
(fig. 304). The residue is composed partly of cyanide of mercury and of
para-cyanogen; it constitutes a very poisonous substance, which should
be immediately thrown away or burned. It can be easily powdered into
dust in the fingers. During the decomposition of the sulpho-cyanide of
mercury, quantities of sulphurous acid are thrown off, and it is to be
regretted that Pharaoh’s serpent should herald his appearance by such a
disagreeable, suffocating odour.
[Illustration: Fig. 303.—Iodide of cyanogen.]
[Illustration: Fig. 304.—Pharaoh’s serpent.]
After these few preliminary experiments, we will endeavour to show
the interest afforded by the study of chemistry in relation to the
commonest substances of every-day life. We will first consider the
nature of a few pinches of salt. We know that kitchen salt, or sea
salt, is white or greyish, according to its degree of purity; that it
has a peculiar flavour, is soluble in water, and makes a peculiar
crackling when thrown in the fire. But though its principal physical
properties may be familiar enough, many people are entirely ignorant of
its chemical nature and elementary composition. Kitchen salt contains
a metal, combined with a gas possessing a very suffocating odour; the
metal is _sodium_, the gas is _chlorine_. The scientific name for the
substance is _chloride of sodium_ (salt).[19] The metal contained in
common salt in no way resembles ordinary metals; it is white like
silver, but tarnishes immediately in contact with air, and unites with
oxygen, thus transforming itself into _oxide of sodium_. To preserve
this singular metal it is necessary to protect it from the action of
the atmosphere, and to keep it in a bottle containing oil of naptha.
Sodium is soft, and it is possible with a pair of scissors to cut it
like a ball of soft bread that has been kneaded in the hand. It is
lighter than water, and when placed in a basin of water floats on the
top like a piece of cork; only it is disturbed, and takes the form of
a small brilliant sphere; great effervescence is also produced as it
floats along, for it reduces the water to a common temperature by its
contact. By degrees the small metallic ball disappears from view, after
blazing into flame (fig. 305).
[Illustration: Fig. 305.—Combustion of sodium in water.]
This remarkable experiment is very easy to carry out, and sodium is now
easily procured at any shop where chemicals are sold. The combustion
of sodium in water can be explained in a very simple manner. Water, as
we know, is composed of hydrogen and oxygen. Sodium, by reason of its
great affinity for the latter gas, combines with it, and forms a very
soluble oxide; the hydrogen is released and thrown off, as we shall
perceive by placing a lighted match in the jar, when the combustible
gas ignites.
Oxide of sodium has a great affinity for water; it combines with it,
and absorbs it in great quantities. It is a solid, white substance,
which burns and cauterizes the skin; it is also _alkaline_, and brings
back the blue colour to litmus paper that has been reddened by acids.
Sodium combines easily also with chlorine. If plunged into a jar
containing this gas it is transformed into a substance, which is sea
salt. If the chlorine is in excess a part of the gas remains free, for
simple substances do not mingle in undetermined ratios; they combine,
on the contrary, in very definite proportions, and 35·5 gr. of dry
chlorine always unites with the same quantity of soda equal to 23
grams. A gram of kitchen salt is formed, therefore, of 0·606 gr. of
chlorine, and 0·394 gr. of sodium. Besides sea salt, there are a number
of different salts which may be made the object of curious experiments.
We know that caustic soda, or oxide of sodium, is an alkaline product
possessing very powerful properties; it burns the skin, and destroys
organic substances.
Sulphuric acid is endowed with no less powerful properties; if a little
is dropped on the hand it produces great pain and a sense of burning; a
piece of wood plunged into this acid is almost immediately carbonized.
If we mix forty-nine grams of sulphuric acid and thirty-one grams of
caustic soda a very intense reaction is produced, accompanied by a
considerable elevation of temperature; after it has cooled we have
a substance which can be handled with impunity; the acid and alkali
have combined, and their properties have been reciprocally destroyed.
They have now originated a _salt_ which is _sulphate of soda_. This
substance exercises no influence on litmus paper, and resembles in no
way the substances from which it originated.
There are an infinite number of salts which result in like manner
from the combination of an acid with an alkali or _base_. Some, such
as sulphate of copper, or chromate of potash, are coloured; others,
like sulphate of soda, are colourless. The last-mentioned salt, with a
number of others, will take a crystalline form; if dissolved in boiling
water, and the solution left to stand, we shall perceive a deposit of
transparent prisms of very remarkable appearance. This was discovered
by Glauber, and was formerly called _Glauber’s salts_.
[Illustration: Fig. 306.—Preparation of a solution saturated with
sulphate of soda.]
Sulphate of soda is very soluble in water, and at a temperature of
thirty-three (Centigr.) water can dissolve it in the greatest degree.
If we pour a layer of oil on a solution saturated with Glauber’s
salts, and let it stand, it will not produce crystals; but if we
thrust a glass rod through the oil into contact with the solution,
crystallization will be instantaneous. This singular phenomenon becomes
even more striking when we put the warm concentrated solution into a
slender glass tube, A B, which we close after having driven out the
air by the bubbling of the liquid (fig. 306). When the tube has been
closed, the crystals of sulphate of soda will not form, even with the
temperature at zero; nevertheless the salts, being less soluble cold
than hot, are found in the fluid in a proportion ten times larger than
they would contain under ordinary conditions. If the end of the tube be
broken the salt will crystallize immediately. We will describe another
experiment, but little known and very remarkable, which exhibits in a
striking manner the process of instantaneous crystallizations. Let one
hundred and fifty parts of hyposulphite of soda be dissolved in fifteen
parts of water, and the solution slowly poured into a test-glass,
previously warmed by means of boiling water, until the vessel is about
half-full. One hundred parts of acetate of soda is then dissolved in
fifteen parts of water, and poured slowly into the first solution,
so that they form two layers perfectly distinct from each other. The
two solutions are then covered with a little boiling water, which,
however is not represented in our illustration. After it has been left
to stand and cool slowly, we have two solutions of hyposulphite of
soda and acetate of soda superposed on each other. A thread, at the
end of which is fixed a small crystal of hyposulphite of soda, is then
lowered into the test-glass; the crystal passes through the solution of
acetate without disturbing it, but it has scarcely reached the lower
solution of hyposulphite than the salt crystallizes instantaneously.
(_See_ the test-glass on the left of fig. 307.) We then lower into the
upper solution a crystal of acetate of soda, suspended from another
thread. This salt then crystallizes also. (_See_ experiment glass on
the right of fig. 307.) This very successful experiment is one of the
most remarkable belonging to the subject of instantaneous crystals.
The successive appearance of crystals of hyposulphite of soda, which
take the form of large, rhomboidal prisms, terminating at the two
extremities with an oblique surface, and the crystals of acetate of
soda, which have the appearance of rhomboidal, oblique prisms, cannot
fail to strike the attention and excite the interest of those who are
not initiated into these kinds of experiments.
[Illustration: Fig. 307.—Experiment of instantaneous crystallization.]
Another remarkable instantaneous crystallization is that of alum. If we
leave standing a solution of this salt it gradually cools, at the same
time becoming limpid and clear. When it is perfectly cold, if we plunge
into it a small octahedral crystal of alum suspended from a thread, we
perceive that crystallization instantly commences on the surface of the
small crystal; it rapidly and perceptibly increases in size, until it
nearly fills the whole jar.
COMMON METALS AND PRECIOUS METALS.
How many invalids have swallowed _magnesia_ without suspecting that
this powder contains a metal nearly as white as silver, and is
malleable, and capable of burning with so intense a light that it
rivals even the electric light in brilliancy! If any of our readers
desire to prepare magnesium themselves it can be done in the following
manner:—Some white magnesia must be obtained from the chemist, and
after having been calcined, must be submitted to the influence of
hydrochloric acid and hydrochlorate of ammonia. A clear solution will
thus be obtained, which by means of evaporation under the influence
of heat, furnishes a double chloride, hydrated and crystallised. This
chloride, if heated to redness in an earthenware crucible, leaves as
a residue a nacreous product, composed of micaceous, white scales,
chloride of anhydrous magnesium.
[Illustration: Fig. 308.—Group of alum crystals.]
If six hundred grams of this chloride of magnesium are mixed with one
hundred grams of chloride of sodium, or kitchen salt, and the same
quantity of fluoride of calcium and metallic sodium in small fragments,
and the mixture is put into an earthenware crucible made red-hot, and
heated for a quarter of an hour under a closed lid, we shall find on
pouring out the fluid on to a handful of earth, that we have obtained
instead of scoria, forty-five grams of metallic magnesium. The metal
thus obtained is impure, and to remove all foreign substances it
must be heated in a charcoal tube, through which passes a current of
hydrogen.
Magnesium is now produced in great abundance, and is very inexpensive.
It is a metal endowed with a great affinity for oxygen, and it is
only necessary to thrust it into the flame of a candle to produce
combustion; it burns with a brightness that the eye can scarcely
tolerate, and is transformed into a white powder—oxide of magnesium,
or magnesia. Combustion is still more active in oxygen, and powder
of magnesium placed in a jar filled with this gas produces a perfect
shower of fire of very beautiful effect. To give an idea of the
lighting power of magnesium, we may add that a wire of this metal,
which is 29/100 of a millimetre in diameter, produces by combustion a
light equal to that of seventy-four candles.
[Illustration: Fig. 309.—Calcined alum.]
[Illustration: Fig. 310.—Preparation of metallic iron.]
The humble earth of the fields—the clay which is used in our potteries,
also contains aluminium, that brilliant metal which is as malleable
as silver, and unspoilable as gold. When clay is submitted to the
influence of sulphuric acid and chloride of potassium, we obtain alum,
which is a sulphate of alumina and potash. Alum is a colourless salt,
which crystallizes on the surface of water in beautiful octahedrons of
striking regularity. Fig. 308 represents a group of alum crystals. This
salt is much used in the colouring of fabrics; it is also used for the
sizing of papers, and the clarification of tallow. Doctors also use
it as an astringent and caustic substance. When alum is submitted to
the action of heat in an earthenware crucible, it loses the water of
crystallization which it contains, and expands in a singular manner,
overflowing from the jar in which it is calcined (fig. 309).
Iron, the most important of common metals, rapidly unites with oxygen,
and, as we know, when a piece of this metal is exposed to the influence
of damp air, it becomes covered with a reddish substance. In the
well-known experiment of the formation of rust, the iron gradually
oxidises without its temperature rising, but this combination of iron
with oxygen is effected much more rapidly under the influence of heat.
If, for example, we redden at the fire a nail attached to a wire, and
give it a movement of rotation as of a sling, we see flashing out
from the metal a thousand bright sparks due to the combination of
iron with oxygen, and the formation of an oxide. Particles of iron
burn spontaneously in contact with air, and this property for many
centuries has been utilized in striking a tinder-box; that is to say,
in separating, by striking a flint, small particles of iron, which
ignite under the influence of the heat produced by the friction. We can
prepare iron in such atoms that it ignites at an ordinary temperature
by simple contact with the air. To bring it to this state of extreme
tenuity, we reduce its oxalate by hydrogen. We prepare an apparatus for
hydrogen as shown in fig. 310, and the gas produced at A is
passed through a desiccative tube, B, and finally reaches a
glass receptacle, C, in which some oxalate of iron is placed.
The latter salt, under the combined influence of hydrogen and heat,
is reduced to metallic iron, which assumes the appearance of a fine
black powder. When the experiment is completed the glass vessel is
closed, and the iron, thus protected from contact with the air, can be
preserved indefinitely; but if it is exposed to the air by breaking
off the end of the receptacle (fig. 311), it ignites immediately,
producing a shower of fire of very beautiful effect. Iron thus prepared
is known under the name of _pyrophoric iron_. Iron is acted upon in
a very powerful manner by most acids. If some nitric acid is poured
on iron nails, a stream of red, nitrous vapour is let loose, and the
oxidised iron is dissolved in the liquid to the condition of nitrate
of iron. This experiment is very easy to perform, and it gives an idea
of the energy of certain chemical actions. We have endeavoured to
represent its appearance in fig. 312. Fuming nitric acid does not act
on iron, and prevents it being attacked by ordinary nitric acid. This
property has given rise to a very remarkable experiment on passive
iron. It consists in placing some nails in a glass, into which some
fuming nitric acid is poured, which produces no result; the fuming acid
is then taken out, and is replaced by ordinary nitric acid, which no
longer acts on the iron rendered passive by the smoking acid. After
this, if the nails are touched by a piece of iron, which has not
undergone the action of nitric acid, they are immediately acted upon,
and a giving off of nitrous vapour is manifested with great energy.
Lead is a very soft metal, and can even be scratched by the nails. It
is also extremely pliable, and so entirely devoid of elasticity that
when bent it has no tendency whatever to return to its primitive form.
Lead is heavy, and has a density represented by 11·4; that is to say,
the weight of a quart of water being one kilogram, that of the same
volume of lead is 11·400 k.
[Illustration: Fig. 311.—Pyrophoric iron.]
Fig. 313 represents cylindrical bars of the best known metals, all
weighing the same, showing their comparative density.
[Illustration: Fig. 312.—Iron and nitric acid.]
Lead, like tin, is capable of taking a beautiful crystalline form
when placed in solution by a metal that is less oxydisable. The
crystallization of lead, represented in fig. 314, is designated by the
name of the _Tree of Saturn_. This is how the experiment is produced:
Thirty grams of acetate of lead are dissolved in a quart of water, and
the solution is poured into a vase of a spherical shape. A stopper for
this vase is made out of a piece of zinc, to which five or six separate
brass wires are attached; these are plunged into the fluid, and we
see the wires become immediately covered with brilliant crystallized
spangles of lead, which continue increasing in size.
The alchemists, who were familiar with this experiment, believed that
it consisted in a transformation of copper into lead, while in reality
it only consists in the substitution of one metal for another. The
copper is dissolved in the liquid, and is replaced by the lead, but no
metamorphosis is brought about. We may vary at will the form of the
vase or the arrangement of the wire; thus it is easy to form letters,
numbers, or figures, by the crystallization of brilliant spangles.
[Illustration: Platinum Density 21.50
Gold D. 19.25
Mercury D. 13.56
Lead D. 11.35
Silver D. 10.47
Bismuth D. 9.82
Copper D. 8.78
Nickel D. 8.27
Tin D. 7.29
Iron D. 7.20
Zinc D. 6.86
Aluminium D. 2.56
Magnesium D. 1.43
Sodium D. 0.97
Fig. 313.—Representation of bars of metal, all of the same weight.]
Copper, when it is pure, has a characteristic red colour, which
prevents it being confounded with any other metal; it dissolves easily
in nitric acid, and with considerable effervescence, giving off vapour
very abundantly. This property has been put to good use in engraving
with aqua fortis. A copper plate is covered with a layer of varnish,
and when it is dry some strokes are made on it by means of a graver;
if nitric acid is poured on the plate when thus prepared, the copper
is only acted on in the parts that have been exposed by the point of
steel. By afterwards removing the varnish, we have an engraved plate,
which will serve for printing purposes.
[Illustration: Fig. 314.—Tree of Saturn.]
Among experiments that may be attempted with common metals, we may
mention that in which salts of tin are employed. Tin has a great
tendency to assume a crystalline form, and it will be easy to show
this property by an interesting experiment. A concentrated solution
of proto-chloride of tin, prepared by dissolving some metallic tin in
hydrochloric acid, is placed in a test glass; then a rod of tin is
introduced, as shown in fig. 315. Some water is next slowly poured on
the rod, so that it gradually trickles down, and prevents the mingling
of the proto-chloride of tin. The vessel is then left to stand,
and we soon see brilliant crystals starting out from the rod. This
crystallization is not effected in the water; it is explained by an
electric influence, into the details of which we cannot enter without
overstepping our limits; it is known as “Jupiter’s Tree.” It is well
known that alchemists, with their strange system of nomenclature,
believed there was a certain mysterious relation between the seven
metals then known and the seven planets; each metal was dedicated to
a planet; tin was called Jupiter; silver, Luna; gold, Sol; lead,
Saturn; iron, Mars; quicksilver, Mercury; and copper, Venus. The
crystallization of tin may be recognised also by rubbing a piece of
this metal with hydrochloric acid; the fragments thus rubbed off
exhibit specimens of branching crystals similar to the hoar-frost which
we see in severe winter weather. If we bend a rod of tin in our hands
the crystals break, with a peculiar rustling sound.
When speaking of precious metals, we may call to mind that the
alchemists considered gold as the king of metals, and the other
valuable ones as noble metals. This definition is erroneous, if we
look upon the useful as the most precious; for, in that case, iron and
copper would be placed in the first rank. If gold were found abundantly
on the surface of the soil, and iron was extremely rare, we should seek
most eagerly for this useful metal, and should despise the former,
with which we can neither make a ploughshare nor any other implement
of industry. Nevertheless, the scarcity of gold, its beautiful yellow
colour, and its unalterability when in contact with air, combine to
place it in the first rank in the list of precious metals. Gold is very
heavy; its density is represented by the figure 19·5. It is the most
malleable and the most ductile of metals, and can be reduced by beating
to such thin sheets that ten thousand can be laid, one over the other,
to obtain the thickness of a millimetre. With a grain of gold a thread
may be manufactured extending a league in length, and so fine that it
resembles a spider’s web. When gold is beaten into thin sheets it is
no longer opaque; if it is fastened, by means of a solution of gum, on
to a sheet of glass, the light passes right through it, and presents
a very perceptible green shade. Gold is sometimes found scattered in
sand, in a condition of impalpable dust, and, in certain localities,
in irregular lumps of varying size, called nuggets. Gold is the least
alterable of the metals, and can be exposed, indefinitely, to the
contact of humid atmosphere without oxidizing. It is not acted on by
the most powerful acids, and only dissolves in a mixture of nitric acid
and hydrochloric acid. We can prove that gold resists the influence of
acids by the following operation:—
Some gold-leaf is placed in two small phials, the first containing
hydrochloric acid, and the second nitric acid. The two vessels are
warmed on the stove, and whatever the duration of the ebullition of
the acids, the gold-leaf remains intact, and completely resists their
action. If we then empty the contents of one phial into the other,
the hydrochloric and nitric acids are mixed, and we see the gold-leaf
immediately disappear, easily dissolved by the action of the liquid
(_aqua regia_). Gold also changes when in contact with mercury; this is
proved by suspending some gold-leaf above the surface of this liquid
(fig. 316); it quickly changes, and unites with the fumes of the
mercury, becoming of a greyish colour.
Silver is more easily affected than gold, and though so white when
fused, tarnishes rapidly in contact with air. It does not oxidize, but
sulphurizes under the influence of hydro-sulphuric emanations. Silver
does not combine directly with the oxygen of the atmosphere; but under
certain conditions it can dissolve great quantities of this gas. If it
is fused in a small bone cupel, in contact with the air, and left to
cool quickly, it expands in a remarkable manner, and gives off oxygen.
[Illustration: Fig. 315.—Jupiter’s Tree.]
Nitric acid dissolves silver very easily, by causing the formation
of abundant fumes. When the solution evaporates, we perceive white
crystals forming, which are nitrate of silver. This fused nitrate of
silver takes the name of _lunar caustic_, and is employed in medicine.
Nitrate of silver is very poisonous; it possesses the singular property
of turning black under the action of the sun’s rays, and is used in
many curious operations in photography. It is also employed in the
manufacture of dyes for the hair; it is applied to white hair with
gall-nut, and under the influence of the light it turns black, and
gives the hair a very dark shade. Salts of silver in solution with
water have the property of forming a precipitate under the influence of
chlorides, such as sea salt. If a few grains of common salt are thrown
into a solution of nitrate of silver, it forms an abundant precipitate
of chloride of silver, which blackens in the light. This precipitate,
insoluble in nitric acid, dissolves very easily in ammonia.
Platinum, which is the last of the precious metals that we have to
consider, is a greyish-white colour, and like gold is only affected by
a mixture of nitric acid and hydrochloric acid. It is the heaviest of
all the ordinary metals; its density is 21·50. It is very malleable
and ductile, and can be beaten into very thin sheets, and into wires as
slender as wires of gold. Platinum wires have even been made so fine
that the eye can scarcely perceive them; these are known as Wollaston’s
invisible wires. Platinum resists the action of the most intense fire,
and we can only fuse it by means of a blow-pipe and hydro-oxide gas.
Its inalterability and the resistance it opposes to fire render it very
valuable for use in the laboratory. Small crucibles are made of it,
which are used by chemists to calcine their precipitates in analytical
operations, or to bring about reactions under the influence of a high
temperature. Platinum may be reduced to very small particles; it then
takes the form of a black powder. In this pulverulent condition it
absorbs gases with great rapidity, to such an extent that a cubic
centimetre can condense seven hundred and fifty times its own volume
of hydrogen gas. It also condenses oxygen, and in a number of cases
acts as a powerful agent. Platinum is also obtained in porous masses
(“spongy platinum”), which produce phenomena of oxidation.
[Illustration: Fig. 316.—Gold-leaf exposed to the fumes of mercury.]
A very ingenious little lamp may be constructed which lights of
itself without the help of a flame. It contains a bell of glass,
which is filled with hydrogen gas, produced by the action exercised
by a foundation of zinc on acidulated water. If the knob on the upper
part of the apparatus is pressed, the hydrogen escapes, and comes in
contact with a piece of spongy platinum, which, acting by oxidation,
becomes ignited. The flame produced sets fire to a small oil lamp,
which is opposite the jet of gas. This very ingenious lamp is known
under the name of Gay-Lussac’s lamp. Platinum can also produce, by mere
contact, a great number of chemical reactions. Place in a test glass an
explosive mixture formed of two volumes of hydrogen and one volume of
oxygen; in this gas plunge a small piece of spongy platinum, and the
combination of the two bodies will be instantly brought about, making
a violent explosion. Make a small spiral of platinum red-hot in the
flame of a lamp, having suspended it to a card; then plunge it quickly
into a glass containing ether, and you will see the metallic spiral
remain red for some time, while in the air it would cool immediately.
This phenomenon is due to the action of oxidation which the platinum
exercises over the fumes of ether. This curious experiment is known
under the name of _the lamp without a flame_. This remarkable oxidizing
power of platinum, which has not yet been explained, was formerly
designated by the title of _catalytic action_. But a phrase is not a
theory, and it is always preferable to avow one’s ignorance than to
simulate an apparent knowledge. Science is powerful enough to be able
to express her doubts and uncertainties boldly. In observing nature we
find an experience of this, and often meet with facts which may be put
to profit, and become useful in application; nevertheless it is often
the case that the why and the wherefore will for a long time escape the
most penetrating eye and lucid intelligence. It is true the admirable
applications of science strike us with the importance of their results,
and the wonderful inventions they originate; but if they turn to
account the observed facts of nature, what do they teach us as to the
first cause of all things, the _wherefore_ of nature?—Almost nothing.
We must humbly confess our powerlessness, and say with d’Alembert: “The
encyclopædia is very abundant, but what of that if it discourses of
what we do not understand?”
[Illustration: Fig. 317.—Discolouration of periwinkles by sulphuric
acid.]
ARTIFICIAL COLOURING OF FLOWERS.
In a course of chemistry, the action exercised by sulphurous acid
on coloured vegetable matter is proved by exposing violets to the
influence of this gas, which whitens them instantaneously. Sulphurous
acid, by its dis-oxidating properties, destroys the colour of many
flowers, such as roses, periwinkles, etc. The experiment succeeds very
readily by means of the little apparatus which we give in fig. 317.
We dissolve in a small vessel some sulphur, which ignites in contact
with air, and gives rise, by its combination with oxygen, to sulphurous
acid; the capsule is covered with a conical chimney, made out of a thin
sheet of copper, and at the opening at the top the flowers that are to
be discoloured are placed. The action is very rapid, and a few seconds
only are necessary to render roses, periwinkles, and violets absolutely
white.
[Illustration: Fig. 318.—Experiment for turning columbines a green
colour with ammoniacal ether.]
M. Filpol, a distinguished _savant_, has exhibited to the members of
the Scientific Association, Paris, the results which he obtained by
subjecting flowers to the influence of a mixture of sulphuric ether
and some drops of ammonia; he has shown that, under the influence of
this liquid, a great number of violets or roses turn a deep green.
We have recently made on this subject a series of experiments which
we will here describe, and which may be easily attempted by those of
our readers who are interested in the question. Some common ether is
poured into a glass, and to it is added a small quantity of liquid
ammonia (about one-tenth of the volume). The flowers with which it
is desired to experiment are then plunged into the fluid (fig. 318).
A number of flowers, whose natural colour is red or violet, take
instantaneously a bright green tint; these are red geranium, violet,
periwinkle, lilac, red and pink roses, wall-flower, thyme, small
blue campanula, fumeter, myosotis, and heliotrope. Other flowers,
whose colours are not of the same shade, take different tints when
in contact with ammoniacal ether. The upper petal of the violet
sweet-pea becomes dark blue, whilst the lower petal turns a bright
green colour. The streaked carnation becomes brown and bright green.
White flowers generally turn yellow, such as the white poppy, the
variegated snow-dragon, which becomes yellow and dark violet, the
white rose, which takes a straw colour, white columbine, camomile,
syringa, white daisy, potatoe blossom, white julian, honeysuckle, and
white foxglove, which in contact with ammoniacal ether assume more or
less deep shades of yellow. White snap-dragon becomes yellow and dark
orange. Red geranium turns blue in a very remarkable fashion; with
the monkey-flower the ammoniacal ether only affects the red spots,
which turn a brownish green; red snap-dragon turns a beautiful brown;
valerian takes a shade of grey; and the red corn-poppy assumes a dark
violet. Yellow flowers are not changed by ammoniacal ether; buttercups,
marigolds, and yellow snap-dragon preserve their natural colour. Leaves
of a red colour are instantly turned green when placed in contact with
ammoniacal ether. The action of this liquid is so rapid that it is
easy to procure green spots by pouring here and there a drop of the
solution. In like manner violet flowers, such as periwinkles, can be
spotted with white, even without gathering them. We will complete our
remarks on this subject with a description of experiments performed
by M. Gabba in Italy by means of ammonia acting on flowers. M. Gabba
simply used a plate, in which he poured a certain quantity of solution
of ammonia. He placed on the plate a funnel turned upside down, in the
tube of which he arranged the flowers on which he wished to experiment.
He then found that under the influence of the ammonia the blue, violet,
and purple flowers became a beautiful green, red flowers black, and
white yellow, etc.
The most singular changes of colour are shown by flowers which are
composed of different tints, their red streaks turning green, the
white yellow, etc. Another curious example is that of red and white
fuchsias, which, through the action of ammonia, turn yellow, blue, and
green. When flowers have been subjected to these changes of colour,
and afterwards plunged into pure water, they preserve their new tint
for several hours, after which they gradually return to their natural
colour. Another interesting observation, due to M. Gabba, is that
asters, which are naturally inodorous, acquire an agreeable aromatic
odour under the influence of ammonia. Asters of a violet colour become
red when wetted with nitric acid mixed with water. On the other hand,
if these same flowers are enclosed in a wooden box, where they are
exposed to the fumes of hydrochloric acid, they become, in six hours’
time, a beautiful red colour, which they preserve when placed in a dry,
shady place, after having been properly dried. Hydrochloric acid has
the effect of making flowers red that have been rendered green by the
action of ammonia, and also alters their appearance very sensibly. We
may also mention, in conclusion, that ammonia, combined with ether,
acts much more promptly than when employed alone.
PHOSPHORESCENCE.
Artificial flowers are frequently to be seen prepared in a particular
manner, which have the property of becoming phosphorescent in darkness,
when they have been exposed to the action of a ray of light, solar or
electric. These curious chemical objects are connected with some very
interesting phenomena and remarkable experiments but little known at
the present time, to which we will now draw the reader’s attention.
The faculty possessed by certain bodies of emitting light when placed
in certain conditions, is much more general than is usually supposed.
M. Edmond Becquerel, to whom we owe a remarkable work on this subject,
divides the phenomena of phosphorescence into five distinct classes:
1. _Phosphorescence through elevation of temperature._ Among the
substances which exhibit this phenomenon in a high degree we may
mention certain diamonds, coloured varieties of fluoride of calcium,
some minerals; and sulphur, known under the name of artificial
phosphorus, when it has previously been exposed to the action of the
light.
2. _Phosphorescence through mechanical action._ This is to be observed
when we rub certain bodies together, or against a hard substance. If we
rub together two quartz crystals in the dark, we perceive red sparks;
and when pounding chalk or sugar, there is also an emission of sparks.
3. _Phosphorescence through electricity._ This is manifested by the
light accompanying disengagement of electricity, and when gases and
rarefied vapours transmit electric discharges.
4. _Spontaneous Phosphorescence_ is observed, as every one knows,
in connection with several kinds of living creatures,—glow-worms,
noctilucids, etc., and similar phosphorescent effects are produced also
with organic substances, animal or vegetable, before putrefaction sets
in. It is manifested also at the flowering time of certain plants, etc.
5. _Phosphorescence through insolation and the action of light._ “It
consists,” says M. Edmond Becquerel, “in exposing for some instants to
the action of the sun, or to that of rays emanating from a powerful
luminous source, certain mineral or organic substances, which
immediately become luminous, and shine in the dark with a light, the
colour and brilliancy of which depend on their nature and physical
character; the light gradually diminishes in intensity during a period
varying from some seconds to several hours. When these substances are
exposed anew to the action of light, the same effect is reproduced.
The intensity of the light emitted after insolation is always much less
than that of the incidental light.” These phenomena appear to have been
first observed with precious stones; then, in 1604, in calcined Bologna
stone, and later, in a diamond by Boyle, in 1663; in 1675 it was
noticed in Baudoin phosphorus (residuum of the calcination of nitrate
of lime), and more recently still in connection with other substances
which we will mention. The substances most powerfully influenced by
the action of light are sulphates of calcium and barium, sulphate of
strontium, certain kinds of diamonds, and that variety of fluoride of
calcium, which has received the name of _chlorophane_.
[Illustration: Fig. 319.—Artificial flower coated with phosphorescent
powder, exposed to the light of magnesium wire.]
Phosphorescent sulphate of calcium is prepared by calcining in an
earthenware crucible a mixture of flowers of sulphur and carbonate
of lime. But the preparation only succeeds with carbonate of lime of
a particular character. That obtained from the calcination of oyster
shells produces very good results. Three parts of this substance is
mixed with one part of flowers of sulphur, and is made red-hot in a
crucible covered in from contact with the air. The substance thus
obtained gives, after its insolation, a yellow light in the dark. The
shells of oysters, however, are not always pure, and the result is
sometimes not very satisfactory; it is therefore better to make use of
some substance whose composition is more to be relied on.
“When we desire to prepare a phosphorescent sulphate with lime,
or carbonate of lime,” says M. E. Becquerel, “the most suitable
proportions are those which in a hundred parts of the substance are
composed of eighty to a hundred of flowers of sulphur in the first
case, and forty-eight to a hundred in the second, that is, when we
employ the quantity of sulphur which will be necessary for burning with
carbonate of lime to produce a monosulphate.[20] It is necessary to
have regard to the elevation of the temperature in the preparation. By
using lime procured from arragonite, and reducing the temperature below
five hundred degrees for a sufficient time for the reaction between the
sulphur and lime to take place, the excess of sulphur is eliminated,
and we have a feebly luminous mass, of a bluish tint; if this mass is
raised to a temperature of eight hundred or nine hundred degrees, it
will exhibit a very bright light.”
Sulphate of calcium possesses different phosphorescent properties
according to the nature of the salt which has served to produce the
carbonate of lime employed. If we transform marble into nitrate of
lime, by dissolving it in water and nitric acid, and form a precipitate
with carbonate of ammonium, and use the carbonate of lime thus obtained
in the preparation of sulphate of calcium, we have a product which
gives a phosphorescence of a violet-red colour. If the carbonate
of lime used is obtained from chloride of calcium precipitated by
carbonate of ammonia, the phosphorescence is yellow. If we submit
carbonate of lime, prepared with lime water and carbonic acid, to the
influence of sulphur, we obtain a sulphur giving a phosphorescent
light of very pure violet. Carbonate of lime obtained by forming a
precipitate of crystallized chloride of calcium with different alkaline
carbonates also gives satisfactory results.
Luminous sulphates of strontium may be obtained, like those of calcium,
by the action of sulphur on strontia or the carbonate of this base, by
the reduction of sulphates of strontia with charcoal. Blue and green
shades are the most common. Sulphates of barium also present very
remarkable phenomena of phosphorescence; but to obtain very luminous
intensity a higher temperature is needed than with the other substances
mentioned, and we have the same result when we reduce native sulphate
of baryta with charcoal; that is to say, when the reaction takes
place which produces the phosphorus formerly known as _phosphorus of
Bologna_. Preparations obtained from baryta have a phosphorescence
varying from orange-red to green.
The preparation of such substances as we have just enumerated afford an
easy explanation of the method of manufacturing the luminous flowers
which we described at the commencement of this chapter. We obtain
some artificial flowers, cover them with some liquid gum, sprinkle
with phosphorescent sulphur, and let them dry. The pulverulent matter
then adheres to them securely, and it is only necessary to expose the
flowers thus prepared to the light of the sun, or the rays emanating
from magnesium wire in a state of combustion (fig. 319), to produce
immediate phosphorescent effects. If taken into a dark room (fig. 320)
they shine with great brilliancy, and give off very exquisite coloured
rays. Phosphorescent sulphates are used also in tracing names or
designs on a paper surface, etc., and it can easily be conceived that
such experiments may be infinitely varied according to the pleasure of
the experimenter.
[Illustration: Fig. 320.—Phosphorescent flower emitting light in a dark
room.]
But let us ask ourselves if these substances are not capable of
being put to more serious uses, and of being classed among useful
products. To this we can reply very decidedly in the affirmative. With
phosphorescent matter we can obtain luminous faces for clocks placed
in dark, obscure spots, and it is not impossible to use it for making
sign-boards for shops, or numbers of houses, which can be lit up at
night. Professor Norton even goes so far as to propose in the “Journal
of the Franklin Institute,” not only coating the walls of rooms with
these phosphorescent substances, but also the fronts of houses, when he
considers it would be possible to do away entirely with street lights,
the house-fronts absorbing sufficient light during the day to remain
luminous the whole of the night.
CHEMISTRY APPLIED TO SLEIGHT OF HAND.
While physics has provided the species of entertainment called “sleight
of hand” with a number of interesting effects, chemistry has only
offered it very feeble contributions. Robert Houdin formerly made use
of electricity to move the hands of his magic clock, and the electric
magnet in making an iron box so heavy instantaneously that no one
could lift it. Robin has made use of optics to produce the curious
spectacle of the decapitated man, spectres, etc. Those persons who are
fond of this kind of amusement may, however, borrow from chemistry
some original experiments, which can be easily undertaken, and I will
conclude this chapter by describing a juggling feat which I have seen
recently executed before a numerous audience by a very clever conjuror.
[Illustration: Fig. 321.—Amusing experiment in chemistry.]
The operator took a glass that was perfectly transparent, and placed
it on a table, announcing that he should cover the glass with a
saucer, and then, retiring to some distance, would fill it with the
smoke from a cigarette. And this he carried out exactly, standing
smoking his cigarette in the background, while the glass, as though
by enchantment, slowly filled with the fumes of the smoke. This trick
is easily accomplished. It is only necessary to pour previously into
the glass two or three drops of hydrochloric acid, and to moisten the
bottom of the saucer with a few drops of ammonia. These two liquids
are unperceived by the spectators, but as soon as the saucer is placed
over the glass, they unite in forming white fumes of hydrochlorate of
ammonia, which bear a complete resemblance to the smoke of tobacco.
This experiment excited the greatest astonishment among the spectators
present on the occasion, but understanding something of chemistry
myself, I easily guessed at the solution of the mystery. The same
result is obtained in a course of chemistry in a more simple manner,
and without any attempt at trickery, by placing the opening of a
bottle of ammonia against the opening of another bottle containing
hydrochloric acid.
FOOTNOTES:
[19] It is the same with a number of other common products, such as
clay, sandstone, etc., the composition of which chemistry has revealed.
Argil, or clay, slate, and schist all contain a metal—_aluminium_,
which has become most valuable for industrial purposes. Stones
for building are composed of a metal combined with carbon and
oxygen—_calcium_; sandstone is composed of _silicium_, a metallic body
united with oxygen; and sulphate of magnesia, which enters into the
composition of a purgative drink, also contains a metal—_magnesium_.
[20] These substances must be finely powdered and thoroughly mixed.