XXV. 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.
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.
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).
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.
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.
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.
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.
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.
Fig. 313 represents cylindrical bars of the best known metals, all weighing the same, showing their comparative density.
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.
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.
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.
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.
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?”
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.
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.
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.
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.
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.
Chapter XXVI
CHEMISTRY AND ALCHEMY—CHEMICAL COMBINATIONS—THE ATMOSPHERIC AIR.
We have in the foregoing pages given some experiments, and considered several of the metals, but there are numerous very interesting subjects still remaining; indeed, the number is so great that we can only pick and choose. All people are desirous to hear something of the atmosphere, of water, and the earth; and as we proceed to speak of crystals and minerals, and so on to geology, we shall learn a good deal respecting our globe—its conformation and constituents. But the atmospheric air must be treated of first. This will lead us to speak of oxygen and nitrogen. Water will serve to introduce hydrogen with a few experiments, and thus we shall have covered a good deal of ground on our way towards various other elements in daily use and appreciation. Now let us begin with a few words concerning CHEMISTRY itself.
At the very outset we are obliged to grope in the dark after the origin of this fascinating science. Shem, or “Chem,” the son of Noah, has been credited with its introduction, and, at any rate, magicians were in Egypt in the time of Moses, and the lawgiver is stated by ancient writers to have gained his knowledge from the Egyptians. But we need not pursue that line of argument. In more modern times the search for the Philosopher’s Stone and the Elixir of Life, which respectively turned everything to gold, and bestowed long life upon the fortunate finder, occupied many people, who in their researches no doubt discovered the germs of the popular science of Chemistry in Alchemy, while the pursuit took a firm hold of the popular imagination for centuries; and even now chemistry is the most favoured science, because of its adaptability to all minds, for it holds plain and simple truths for our every-day experience to confirm, while it leads us step by step into the infinite, pleasing us with experiments as we proceed.
Alchemy was practised by numerous quacks in ancient times and the Middle Ages, but all its professors were not quacks. Astrology and alchemy were associated by the Arabians. Geber was a philosopher who devoted himself entirely to alchemy, and who lived in the year 730 A.D. He fancied gold would cure all disease, and he did actually discover corrosive sublimate, nitric acid, and nitrate of silver. To give even a list of the noted alchemists and magicians would fill too much space. Raymond Sully, Paracelsus, Friar Bacon, Albertus Magnus, Thomas Aquinas, Flamel, Bernard of Treves, Doctor Dee, with his assistant Kelly, and in later times Jean Delisle, and Joseph Balsamo (Cagliostro), who was one of the most notorious persons in Europe about one hundred years ago (1765-1789), are names taken at random; and with the older philosophers chemistry was an all-absorbing occupation—not for gold, but knowledge.
The revelation was slow. On the temperature of bodies the old arts of healing were based—for chemistry and medicine were allies. The elements, we read, existed on the supposition “that bodies were hot or cold, dry or moist”; and on this distinction for a long time “was based the practice of medicine.” The doctrine of the “three principles” of existence superseded this,—the principles being salt, mercury, and sulphur. Metals had been regarded as living bodies, gases as souls or spirits. The idea remained that the form of the substance gave it its character. Acid was pointed; sweet things were round.
Chemistry, then, has had a great deal to contend against. From the time of the Egyptians and Chinese, who were evidently acquainted with various processes,—dyeing, etc.,—the science filtered through the alchemists to Beecher and Stahl, and then the principle of affinity—a disposition to combine—was promulgated, supplemented in 1674 by Mayow, by the theory of divorce or analysis. He concluded that where union could be effected, separation was equally possible. In 1718 the first “Table of Affinities” was produced. Affinity had been shown to be elective, for Mayow pointed out that fixed salts chose one acid rather than another. Richter and Dalton made great advances. Before them Hales, Black, Priestley, Scheele, Lavoisier, and numerous others penetrated the mysteries of the science whose history has been pleasantly written by more than one author who we have not been able to consult, and have no space to do more than indicate. In later days Faraday, De la Rive, Roscoe, and many others have rendered chemistry much more popular, while they have added to its treasures. The story of the progress of chemistry would fill a large volume, and we have regretfully to put aside the introduction and pass on.
Before proceeding to investigate the elements, a few words concerning the general terms used in chemistry will be beneficial to the reader. If we look at the list of the elements, pp. 308-9, we shall see various terminations. Some are apparently named from places, some from their characteristics. Metals lately discovered by the spectroscope (and recently) end in ium; some end in “ine,” some in “on.” As far as possible in late years a certain system of nomenclature has been adhered to, but the old popular names have not been interfered with.
When elements combine together in certain proportions of each they receive certain names. The following table will explain the terms used; for instance, we find that—
Compounds of Oxygen are termed Oxides, as oxide of copper. ” Hydrogen ” Hydrides, as hydride of potassium. ” Chlorine ” Chlorides, as chloride of sodium. ” Nitrogen ” Nitrides, as nitride of boron. ” Bromine ” Bromides, as bromide of potassium. ” Iodine ” Iodides, as iodide of potassium. ” Sulphur ” Sulphides, or} Sulphurets,} as sulphuret of lead. ” Selenium ” Selenides, as selenide of mercury. ” Carbon ” Carbides, or} as carbide of nitrogen, Carburets,} and so on.
The above examples refer to the union in single proportion of each, and are called Binary Compounds. When more than one atom of each element exists in different proportions we have different terms to express such union. If one atom of oxygen be in the compound it is called a “monoxide” or “protoxide”; two atoms of oxygen in combination is termed “dioxide” or “binoxide”; three, “trioxide,” or “tritoxide”; four is the “tetroxide” or “per-oxide,” etc. When more than one atom, but not two atoms is involved, we speak of the sesqui-oxide (one-and-a-half),—“oxide” being interchangeable for “sulphide” or “chloride,” according to the element.
There are other distinctions adopted when metals form two series of combinations, such as ous and ic, which apply, as will be seen, to acids. Sulphur_ic_ and sulphur_ous_ acids, nitr_ic_ and nitr_ous_ acid are familiar examples. In these cases we shall find that in the acids ending in “ous” oxygen is present in less quantity than in the acids ending in ic. The symbolic form will prove this directly, the number of atoms of oxygen being written below,
Sulphurous Acid = H_{2} SO_{3}. Nitrous Acid = HNO_{2}. Sulphuric Acid = H_{2}SO_{4}. Nitric Acid = HNO_{3}.
Whenever a stronger compound of oxygen is discovered than that denominated by ic, chemists adopt the plan of dubbing it the per (ὑπέρ over), as per-chloric acid, which possesses four atoms of oxygen (HClO_{4}), chloric acid being HClO_{3}. The opposite Greek term, ὑπὸ (hupo, below), is used for an acid with less than two atoms of oxygen, and in books is written “hypo”-chlorous (for instance). Care has been taken to distinguish between the higher and lower; for “hyper” is used in English to denote excess, as hyper-critical; and hypo might to a reader unacquainted with the derivation convey just the opposite meaning to what is intended.
While speaking of these terminations we may show how these distinctive endings are carried out. We shall find, if we pursue the subject, that when we have a salt of any acid ending in ic the salt terminates in “ate.” Similarly the salts of acids ending in ous, end in “ite.” To continue the same example we have—
Sulphur_ous_ Acid, which forms salts called Sulph_ites_. Sulphur_ic_ Acid, ” ” Sulph_ates_.
Besides these are sulph_ides_, which are results of the unions or compounds of elementary bodies. Sulph_ites_ are more complicated unions of the compounds. Sulphates are the salts formed by the union of sulphuric acid with bases. Sulphides or sulphurets are compounds in which sulphur forms the electro-negative element, and sulphites are salts formed by the union of sulphur_ous_ acids with bases, or by their action upon them.
The symbolical nomenclature of the chemist is worse than Greek to the uninitiated. We frequently see in so-called popular chemical books a number of hieroglyphics and combinations of letters with figures very difficult to decipher, much less to interpret. These symbols take the place of the names of the chemical compounds. Thus water is made up of oxygen and hydrogen in certain proportions; that is, two of hydrogen to one of oxygen. The symbolic reading is simple, H_{2}O, = the oxide of hydrogen. Potassium again mingles with oxygen. Potassium is K in our list; KO is oxide of potassium (potash). Let us look into this a little closer.
The union of one particle of a simple body with a particle of another simple body can be easily understood; but, as we have seen, it is possible to have substances consisting of four or five different particles, though the greater number of chemical combinations consist of two or three dissimilar ones. In the diagram (fig. 322) we have some possible combinations.
In these combinations we may have one particle of a in combination with one, two, three, four, or five of b, and many particles of a can unite with various molecules of b. Suppose we have oxygen and sulphur compounds as follows:—
Thus there are three different compounds of these two elements—SO, SO_{2}, SO_{3} (without water).
A compound body may combine with another compound body, and this makes a complicated compound. Suppose we have a mixture of sulphuric acid and potash. We have a sulphate of potassium (K_{2}SO_{4}) and combinations of these combinations may likewise be formed. We must read these symbols by the light of the combining weights given in the table, and then we shall find the weight of oxygen or other elements in combination. Thus when we see a certain symbol (Hg.S for instance), we understand that they form a compound including so many parts of mercury and so many of sulphur, which is known as vermilion. Hg.O is oxide of mercury, and by reference to the table of Atomic Weights, we find mercury is Hg., and its combining weight is 200; while oxygen is O, and its weight is 16. Thus we see at once how much of each element is contained in oxide of mercury, and this proportion never varies; there must be 200 of one and 16 of the other, by weight, to produce the oxide. So if the oxygen has to be separated from it, the sum of 216 parts must be taken to procure the 16 parts of oxygen. When we see, as above, O_{2} or O_{3}, we know that the weight must be calculated twice or three times, O being 16; O_{2} is therefore 32 parts by weight. So when we have found what the compounds consist of, we can write them symbolically with ease.
COMPOSITION OF THE ATMOSPHERIC AIR.
We have already communicated a variety of facts concerning the air. We have seen that it possesses pressure and weight. We call the gaseous envelope of the earth the atmosphere, and we are justified in concluding that other planets possess an atmosphere also, though of a different nature to ours. We have seen how easy it is to weigh the air, but we may repeat the experiment. (See illustration, fig. 45, page 50.) We shall find that a perfectly empty glass globe will balance the weights in the scale-pan; admit the air, and the glass globe will sink. So air possesses weight. We have mentioned the Magdeburg hemispheres, the barometer, the air-pump, and the height and the pressure of the atmosphere have been indicated. The density of the atmosphere decreases as we ascend; for the first seven miles the density diminishes one-fourth that of the air at the sea-level, and so on for every succeeding seven.
In consequence of the equal, if enormous, pressure exercised in every direction, we do not perceive the inconvenience, but if the air were removed from inside of a drum, the parchment would quickly collapse. We feel the air when we move rapidly. We breathe the air, and that statement brings us to consider the composition of the atmosphere, which, chemically speaking, may vary a little (as compared with the whole mass) in consequence of changes which are continually taking place, but to all intents and purposes the air is composed as follows, in 100 parts:
Nitrogen 79 parts. Oxygen 20 ” Carbonic Acid .04 ”
with minute quantities of other ingredients, such as ammonia, iodine, carbonetted hydrogen, hydrochloric acid, sulphuretted hydrogen, nitric acid, carbonic oxide, and dust particles, as visible in the sunbeams, added.
The true composition of the atmosphere was not known till Lavoisier demonstrated that it consisted of two gases, one of which was the vital fluid, or oxygen, discovered by Priestley. To the other gas Lavoisier gave the name of Azote,—an enemy of life,—because it caused death if inhaled alone. The carbonic acid in the air varies very much, and in close, heated, and crowded rooms increases to a large quantity, which causes lassitude and headache.
We can easily prove the existence of carbonic acid gas as exhaled from the lungs. Suppose we take a glass and fill it partly with clear lime-water; breathe through a glass tube into the water in the glass, and very quickly you will perceive that the lime-water is becoming cloudy and turbid. This cloudiness is due to the presence of chalk, which has been produced by the action of the carbonic acid gas in the lime-water. This is a well known and always interesting experiment, because it leads up to the vital question of our existence, and the functions of breathing and living.
A popular writer once wrote a book entitled, “Is Life Worth Living?” and a witty commentator replied to the implied question by saying, “It depends upon the liver.” This was felt to be true by many people who suffer, but the scientific man will go farther, and tell you it depends upon the air you breathe, and on the carbonic acid you can raise to create heat,—animal heat,—which is so essential to our well-being. We are always burning; a furnace is within us, never ceasing to burn without visible combustion. We are generating heat by means of the blood. We know that we inhale air into the lungs, and probably are aware that the air so received parts with the oxygen to renew the blood. The nitrogen dilutes the oxygen, for if we inhaled a less-mixed air we should either be burnt up or become lunatics, as light-headed as when inhaling “laughing-gas.” This beautifully graduated mixture is taken into our bodies, the oxygen renews the blood and gives it its bright red colour; the carbon which exists in all our bodies is cold and dead when not so vivified by oxygen. The carbonic acid given off produces heat, and our bodies are warm. But when the action ceases we become cold, we die away, and cease to live. Man’s life exemplifies a taper burning; the carbon waste is consumed as the wax is, and when the candle burns away—it dies! It is a beautiful study, full of suggestiveness to all who care to study the great facts of Nature, which works by the same means in all matter. We will refer to plants presently, after having proved by experiment the existence of nitrogen in the air.
Rutherford experimented very cruelly upon a bird, which he placed beneath a glass shade, and there let it remain in the carbonic acid exhaled from its lungs, till the oxygen being at length all consumed by the bird, it died. When the atmosphere had been chemically purified by a solution of caustic potash, another bird was introduced, but though it lived for some time, it did not exist so long as the first. Again the air was deprived of the carbonic acid, and a third bird was introduced. The experiment was thus repeated, till at length a bird was placed beneath the receiver, and it perished at once. This is at once a cruel and clumsy method of making an experiment, which can be more pleasantly and satisfactorily practised by burning some substance in the air beneath the glass. Phosphorus, having a great affinity for oxygen, is usually chosen. The experiment can be performed as follows with a taper, but the phosphorus is a better exponent.
Let us take a shallow basin with some water in it, a cork or small plate floating upon the water, and in the plate a piece of phosphorus. We must be careful how we handle phosphorus, for it has a habit, well known, but sometimes forgotten by amateur chemists, of suddenly taking fire. Light this piece of phosphorus,—a small piece will do if the jar be of average “shade” size,—and place the glass over it, as in the illustration (fig. 325). The smoke will quickly spread in the jar, and the entry of air being prevented, because the jar is resting under water, phosphoric acid will be formed, and the oxygen thereby consumed. The water, meanwhile, will rise in the jar, the pressure of the air being removed. The burning phosphorus will soon go out, and when the glass is cool, you will be able to ascertain what is inside the jar. Put a lighted taper underneath, and it will go out. The taper would not go out before the phosphorus was burnt in the glass, and so now we perceive we have azote in the receptacle—that is, nitrogen. The other, the constituent of our atmosphere, carbonic acid, as we have seen, is very injurious to the life of animals, and as every animal breathes it out into the air, what becomes of it? Where does all this enormous volume of carbonic acid, the quantities of this poison which are daily and nightly exhaled, where do they all go to? We may be sure nature has provided for the safe disposal of it all. Not only because we live and move about still,—and of course that is a proof,—but because nature always has a compensating law. Remember nothing is wasted; not even the refuse, poisonous air we get rid of from our lungs. Where does it go?
It goes to nourish the plants and trees and vegetables that we delight to look upon and to eat the fruit of. Thus the vegetable world forms a link between the animals and the minerals. Vegetables obtain food, so to speak, and nourishment from water, ammonia, and carbonic acid, all compound bodies, but inorganic.
Water consists of oxygen and hydrogen, carbonic acid of carbon and oxygen, and ammonia of hydrogen and nitrogen. Water and ammonia are present in the air; so are oxygen and nitrogen. Water falls in the form of rain, dew, etc. So in the atmosphere around us we find nearly every necessary for plant-life; and in the ground, which supplies some metallic oxides for their use, we find the remainder. From the air, then, the plant derives its life.
The vegetable kingdom in turn gives all animals their food. This you will see at a glance is true. Certainly animals live on animals. Man and wilder animals live on the beasts of the field in a measure, but those beasts derive their nourishment from vegetables—the vegetable kingdom. So we live on the vegetable kingdom, and it separates the carbonic acid from the air, and absorbs it. What we do not want it takes. What we want it gives. Vegetables give out oxygen, and we consume it gladly. We throw away carbonic acid, and the plants take it greedily; and thus the atmosphere is retained pure for our use. We can, if desirable, prove that plants absorb carbonic acid and give out oxygen by placing leaves of a plant in water, holding the acid in solution, and let the sun shine upon them. Before long we shall find that the carbonic acid has disappeared, and that oxygen has come into the water.
Carbonic acid is sufficiently heavy to be poured from one vessel to another; and if we have obtained some in a glass, we can extinguish a taper by pouring the invisible gas on to the lighted taper, when it will be immediately extinguished.
From the foregoing observations it will be perceived how very desirable it is that ventilation should be attended to. People close up windows and doors and fireplaces, and go to bed and sleep. In the morning they complain of headache and lassitude; they wonder what is the matter, and why the children are not well. Simply because they have been rebreathing the carbonic acid. Go into a closed railway carriage which is nearly filled (and it is astonishing to us how people can be so foolish as to close every outlet), and you will recoil in disgust. These travellers shut the ventilators and windows “because of the cold.” A very small aperture will ventilate a railway carriage; but a close carriage is sickening and enervating, as these kind of travellers find out by the time they reach their journey’s end. Air was given us to breathe at night as well as by day; and though from man’s acts or omissions there may be circumstances in which “night” air may affect the health, we maintain that air is no more injurious naturally than “day” air. Colder it may be, but any air at night is “night” air, in or out of doors at night; and we are certain that night air in itself never hurt any healthy person. It is not nature’s plan to destroy, but to save. If a person delicate in constitution gets hot, and comes out into a colder atmosphere, and defy nature in that way, he (or she) must take the consequences. But air and ventilation (not draught) are necessaries of health, and to say they injure is to accuse nature falsely. There are many impurities in the air in cities, and in country places sometimes, but such impurities are owing to man’s acts and omissions. With average sanitary arrangements and appliances in a neighbourhood no one need be afraid to breathe fresh air night or day; and while many invalids have, we believe, been retarded in recovery from being kept in a close room, hundreds will be benefited by plenty of fresh air. We should not so insist upon these plain and simple truths were there not so many individuals who think it beneficial to close up every avenue by which air can enter, and who then feel ill and out of spirits, blaming everything but their own short-sightedness for the effect of their own acts. An inch or two of a window may be open at night in a room, as the chimney register should be always fully up in bedrooms. When there are fires the draught supplies fresh air to the room with sufficient rapidity. But many seaside journeys might be avoided if fresh air were insisted on at home.
There is another and an important constituent of the atmosphere called OZONE, which is very superior oxygen, or oxygen in what is termed the “Allotropic” state, and is distantly related to electricity, inasmuch as it can be produced by an electrical discharge. This partly accounts for the freshness in the air after a thunderstorm, for we are all conscious that the storm has “cleared the air.” The fresh, crisp ozone in the atmosphere is evident. Ozone differs from oxygen in possessing taste and smell, and it is heavier by one-half than the oxygen gas. There is a good deal of ozone in the sea breeze, and we can, though not infallibly, detect its presence by test-paper prepared with iodide of potassium, which, when ozone is present, will turn blue. We have still something to learn about ozone, which may be considered as “condensed oxygen.”
We have frequently mentioned “combustion,” and as under ordinary circumstances such effects cannot take place without atmospheric air, we will consider it. Combustion is chemical action accompanied by light and heat. Chemical union is always attended by the development of heat, not always by light, because the union varies in intensity and quickness. But when a candle is burning we can study all the interesting phenomena of combustion. We have already spoken of HEAT and LIGHT, so we need only refer the readers to those subjects in the former parts of this volume. Heat is referable to chemical action, and varies according to the energy of union. Heat is always present, remember, in a greater or less degree; and when visible combustion takes place we see light. Invisible combustion goes on in our bodies, and we feel heat; when we get cold we feed the fire by eating, or blow it by exercise and air in our lungs.
We shall speak, however, of combustion now as it affects us in daily life; our fires, our candles, gas, etc., and under these ordinary circumstances hydrogen and carbon are present. (We shall hear more about carbon presently.) These unite with the oxygen to form water and carbonic acid; the water being visible as we first put the cold shade upon the lighted lamp, and the carbonic acid renders the air impure.
In the case of a common candle, or lamp, combustion takes place in the same way. The wick is the intermediary. The oil mounts in the lamp wick, where it is converted into a gas by heat; it then “takes fire,” and gives us light and heat. The candle-flame is just the same with one exception: the burning material is solid, not liquid, though the difference is only apparent, for the wax is melted and goes up as gas. The burning part of the wick has a centre where there is no combustion, and contains carbon. We can prove this by placing a bent tube, as in the illustration (fig. 326), one end in the unburning part of the flame. We shall soon see a dark vapour come over into the receiver. This is combustible, for if we raise the tube without the glass we can light the gas (fig. 327). If we insert the end of the tube into the brilliant portion of the flame we shall perceive a black vapour, which will extinguish the combustion, for it is a mixture of carbonic acid gas and aqueous vapour, in which (fig. 328) particles of carbon are floating.
When we proceed to light our lamps to read or to write by, we find some difficulty in making the wick burn at first. We present to it a lighted taper, and it has no immediate effect. Here we have oil and cotton, two things which would speedily set a warehouse in flames from top to bottom, but we cannot even ignite them, try all we can. Why?—Because we must first obtain a gas, oil will not burn liquid; it must be heated to a gaseous point before it will burn, as all combustion depends upon that,—so flames mount high in air. Now in a candle-flame, as will be seen in the diagram (fig. 331.), there are three portions,—the inner dark core, which consists of unburnt gas; the outer flame, which gives light; and the outside rim of perfect combustion non-luminous. In the centre, A, there is no heat. If we place a piece of gauze wire over the flame at a little distance the flame will not penetrate it. It will remain underneath, because the wire, being of metal, quickly absorbs the heat, and consequently there is no flame. This idea led to the invention of the “safety” lamp by Sir Humphrey Davy, which, although it is not infallible, is the only lamp in general use in mines (figs. 329, 330).
Mines must have light, but there is a gas in mines, a “marsh” gas, which becomes very explosive when it mixes with oxygen. Of course the gas will be harmless till it meets oxygen, but, in its efforts to meet, it explodes the moment the union takes place; instead of burning slowly like a candle it goes off all at once. This gas, called “fire damp,” is carburetted hydrogen, and when it explodes it develops into carbonic acid gas, which suffocates the miners.
XXVII. NON-METALLIC ELEMENTS
OXYGEN—SYMBOL =O=; ATOMIC WEIGHT 16.
Oxygen is certainly the most abundant element in nature. It exists all around us, and the animal and vegetable worlds are dependent upon it. It constitutes in combination about one-half of the crust of the earth, and composes eight-ninths of its weight of water. It is a gas without taste or colour. Oxygen was discovered by Priestley and Scheele, in 1774, independently of each other.
Oxygen can be procured from the oxides of the metals, particularly from gold, silver, and platinum. The noble metals are reducible from their oxides by heat, and this fact assists us at once. If we heat chlorate of potash, mixed with binoxide of manganese, in a retort in a furnace, the gas will be given off. There are many other ways of obtaining oxygen, and we illustrate two (figs. 333, 335).
The red oxide of mercury will very readily evolve oxygen, and if we heat a small quantity of the compound in a retort as per illustration (fig. 333) we shall get the gas. In a basin of water we place a tube test-glass, and the gas from the retort will pass over and collect in the test tube, driving out the water.
The other method mentioned above,—viz., by heating chlorate of potash, etc., in a furnace, is shown in the following illustration. Oxygen, as we have said, is a colourless and inodorous gas, and for a long time it could not be obtained in any other form; but lately both oxygen and hydrogen have been liquified under tremendous pressure at a very low temperature. Oxygen causes any red-hot substance plunged into it to burn brightly; a match will readily inflame if a spark be remaining, while phosphorus is exceedingly brilliant, and these appearances, with many others equally striking, are caused by the affinity for those substances possessed by the gas. Combustion is merely oxidation, just as the process of rusting is, only in the latter case the action is so slow that no sensible heat is produced. But when an aggregate of slowly oxidising masses are heaped together, heat is generated, and at length bursts into flame. This phenomenon is called “spontaneous combustion.” Cases have been known in which the gases developed in the human body by the abuse of alcoholic drinks have ended fatally; in like manner the body being completely charred. (Combustion must not be confounded with ignition, as in the electric light.) Oxygen then, we see, is a great supporter of combustion, though not a combustible itself as coal is. When the chemical union of oxygen with another substance is very rapid an explosion takes place.
Oxidation occurs in various ways. Besides those already mentioned, all verdigris produced on copper, all decays of whatever kind, disintegration, and respiration, are the effects of oxygen. The following experiment for the extraction of oxygen directly from the air was made by M. Boussingault, who passed the gas upon a substance at a certain temperature, and released it at a higher. The illustration on page 351 will show the way in which the experiment was performed.
Boussingault permitted a thin stream of water to flow into a large empty flask, and by this water the air was gradually driven out into a flask containing chloride of calcium and sulphuric acid, which effectually dried it. This dry air then passed into a large tube inside the reverberatory furnace, in which tube were pieces of caustic baryta. Heated to a dull redness this absorbs oxygen, and when the heat is increased to a bright red the superabundant gas is given off. Thus the oxygen was permitted to pass from the furnace-tube into the receiving glass, and so pure oxygen was obtained from the air which had been in the glass bottle at first (fig. 338).
HYDROGEN—SYMBOL =H=; ATOMIC WEIGHT 1.
Hydrogen is abundant in nature, but never free. United with oxygen it forms water, hence its name, “water-former.” It is to Parcelcus that its discovery is due, for he found that oil of vitriol in contact with iron disengaged a gas which was a constituent of water. This gas was subsequently found to be inflammable, but it is to Cavendish that the real explanation of hydrogen is owing. He explained his views in 1766.
Hydrogen is obtained in the manner illustrated in the cut, by means of a furnace, as in fig. 339, or by the bottle method, as per fig. 340. The first method is less convenient than the second. A gun-barrel or fire-proof tube is passed through the furnace, and filled with iron nails or filings; a delivery tube is at the farther end, and a flask of water boiling at the other. The oxygen combines with the iron in the tube, and the hydrogen passes over. The second method is easily arranged. A flask, as in the cut, is provided, and in it some zinc shavings are put. Diluted sulphuric acid is then poured upon the metal. Sulphate of zinc is formed in the flask, and the hydrogen passes off.
Hydrogen being the lightest of all known bodies, its weight is put as 1, and thus we are relatively with it enabled to write down the weights of all the other elements. Hydrogen is fourteen-and-a-half times lighter than atmospheric air, and would do admirably for the inflation of balloons were it not so expensive to procure in such large quantities as would be necessary. Ordinary coal gas, however, contains a great deal of hydrogen, and answers the same purpose.
A very pretty experiment may be made with a bladder full of hydrogen gas. If a tube be fitted to the bladder already provided with a stop-cock, and a basin of ordinary soap-suds be at hand, by dipping the end of the tube in the solution and gently expressing the gas, bubbles will be formed which are of exceeding lightness (fig. 341). They can also be fired with a taper.
Another experiment may be made with hydrogen as follows:—If we permit the gas to escape from the flask, and light it, as in the illustration, and put a glass over it, we shall obtain a musical note, higher or lower, according to the length, breadth, and thickness of the open glass-tube (fig. 342). If a number of different tubes be employed, we can obtain a musical instrument—a gas harmonium.
Hydrogen burns with a blue flame, and is very inflammable. Even water sprinkled upon a fire will increase its fierceness, because the hydrogen burns with great heat, and the oxygen is liberated. Being very light, H can be transferred from one vessel to another if both be held upside down. Some mixtures of H and O are very explosive. The oxyhydrogen blow-pipe is used with a mixture of O and H, which is forcibly blown through a tube and then ignited. The flame thus produced has a most intense heating-power.
A very easy method of producing hydrogen is to put a piece of sodium into an inverted cylinder full of water, standing in a basin of water. The sodium liberates the hydrogen by removing the oxygen from the liquid.
WATER—SYMBOL =H_{2}O=; ATOMIC WEIGHT 18.
At page 59 of this volume we said something about water, and remarked (as we have since perceived by experiment) that “water is composed of oxygen and hydrogen in proportions, by weight, of eight of the former to one of the latter gas; in volume, hydrogen is two to one”; and we saw that “volume and weight were very different things.” This we will do well to bear in mind, and that, to quote Professor Roscoe, “Water is always made up of sixteen parts of oxygen to two parts of hydrogen by weight”; sixteen and two being eighteen, the combining weight of water is eighteen.
We can prove by the Eudiometer that hydrogen when burnt with oxygen forms water; and here we must remark that water is not a mere mechanical mixture of gases, as air is. Water is the product of chemical combination, and as we have before said, is really an oxide of hydrogen, and therefore combustion, or electricity, must be called to our assistance before we can form water, which is the result of an explosion, the mixture meeting with an ignited body—the aqueous vapour being expanded by heat.
The ancients supposed water to be a simple body, but Lavoisier and Cavendish demonstrated its true character. Pure water, at ordinary temperatures, is devoid of taste and smell, and is a transparent, nearly colourless, liquid. When viewed in masses it is blue, as visible in a marked degree in the Rhone and Rhine, at Geneva, and Bâle respectively. Its specific gravity is 1, and it is taken as the standard for Sp. Gravity, as hydrogen is taken as the standard for Atomic Weight. The uses of water and the very important part it plays in the arrangements of nature as a mechanical agent, geology can attest, and meteorology confirm. It composes the greater portions of animals and plants; without water the world would be a desert—a dead planet.
We sometimes speak of “pure” spring water, but such a fluid absolutely pure can scarcely be obtained; and though we can filter water there will always remain some foreign substance or substances in solution. It is well known that the action of water wears away and rounds off hard rocks, and this power of disintegration is supplemented by its strength as a solvent, which is very great. Rain-water is purest in the country as it falls from the clouds. In smoky towns it becomes sooty and dirty. It is owing to the solvent properties of water, therefore, that we have such difficulty in obtaining a pure supply. There is hard water and soft water. The former is derived from the calcareous formations, and contains lime, like the Kent water. This can be ascertained by noticing the incrustations of the vessels wherein the water is boiled. But water rising from hard rocks, such as granite, can do little to disintegrate them at the moment, and therefore the water rises purer. Springs from a great depth are warm, and are known as “thermal springs”; and when they come in contact with carbonic acid and some salts in their passage to the surface, they are known as “mineral waters.” These waters hold in solution salts of lime and magnesia, or carbonates of soda with those of lime and magnesia; salts of iron, and compounds of iodine and bromine are found in the natural mineral waters also, as well as sulphurous impregnations, instances of which will occur to every reader.
We mentioned the Eudiometer just now, and we give an illustration of it. This instrument is used to ascertain the proportions in which the elements of water are composed by synthesis, or a putting together of the constituents of a body to make it up. This is distinguished from analysis, which means separating the compound body into its elements, as we do when we pass the electric current through water.
The Eudiometer consists of a stout glass tube sealed hermetically at one end; two platinum wires are pushed in through the glass just before the end is sealed. The tube is now filled with mercury, and inverted in a bowl of the same metal. Hydrogen, and then oxygen, are admitted through the mercury in the recognised proportions of two to one. By the time the mercury is somewhat more than half displaced, the tube should be held upon a sheet of india-rubber at the bottom of the vessel to keep the metal in the tube, for when the necessary explosion takes place the mercury might also be driven out. A spark from the electrophorus or from a Leyden jar may now be passed through the gases in the tube. The explosion occurs, and water is formed inside. If the mercury be again admitted it will rise nearly to the very top of the tube, driving the bubble up. Thus we find we have formed water from the two gases.
The decomposition of water is easily affected by electricity, and if a little sulphuric acid be added to the water, the experiment will be thereby facilitated. Two wires from a battery should be inserted through a glass filled with the water, and into two test tubes also filled. The wires terminate in platinum strips, and are fastened at the other end to the positive and negative poles of the galvanic battery. The gases will collect in the test tubes, and will be found in proper proportions when the current passes.
So much for water in its liquid state. The solid condition of water (ice) is equally interesting. When we apply heat to water, we get a vapour called “steam”; when we cool water to 32° Fahr., we get a solid mass which weighs just the same as the liquid we have congealed, or the steam we have raised from an equal amount of water. But water expands while in the process of solidification, just as it does when it becomes gaseous, and as we have remarked before, our water-pipes bear full testimony to this scientific fact. When ice forms it has a tendency to crystallize, and some of these ice crystals are, as we see, very beautiful. Snow is only water in a nearly solid form, and the crystals are extremely elegant, appearing more like flowers than congealed water, in tiny six-pointed ice crystals. Many philosophers of late years have written concerning these tiny crystals, which, in common with all crystals, have their own certain form, from which they never depart. Snowflakes are regular six-sided prisms grouped around a centre forming angles of 60° and 120°. There are a number of forms, as will be seen from the accompanying illustrations, and at least ninety-six varieties have been observed. One snowflake, apparently so like all other flakes that fall, can thus be viewed with much interest, and yet, while so very various, snowflakes never get away from their proper hexagonal structure. It has been remarked that snowflakes falling at the same time have generally the same form.
Of the latent heat of ice, etc., we have already spoken in our article upon Heat, and therefore it will be sufficient to state that the latent heat of water is 79 thermal units, because when passing from the liquid to the solid state a certain amount of water absorbs sufficient heat to raise an equal quantity of the liquid 79°. This can be proved by taking a measured quantity (say a pint) of water at 79° and adding ice of the same weight to the water. The mixture will be found to be at zero. Therefore the ice has absorbed or rendered latent 79° of heat which the water possessed. If we melt ice until only a trace of it is left, we shall still find the water as cold as the ice was; all the latent heat is employed in melting the ice. So it will take as much heat to bring a pound of ice at zero to a pound of water at zero, as it would to raise 79 pounds of water 1°. The same law applies to steam.
Water can be distilled in small quantities by an apparatus, as figured in the illustration, and by these means we get rid of all impurities which are inseparable from the liquid otherwise. When it is desirable to distil large quantities of water a larger apparatus is used, called an “Alembic.” The principle is simply to convert the liquid by heat into vapour, then cool it, by condensation, in another vessel.
The evaporation of water, with its effects upon our globe, belong more to the study of Meteorology.
Rain-water is the purest, as we have said, because it goes through the process of distillation by nature. The sun takes it up, by evaporation, into the air, where it is condensed, and falls as rain-water. Water containing carbonate of lime will petrify or harden, as in stalactite caverns. The carbonic acid escapes from the dripping water, the carbonate in solution is deposited as a stalactite, and finally forms pillars in the cave. Sea-water contains many salts; its composition is as follows, according to Dr. Schwertzer, of Brighton:—
Water 964·74372 grains. Chloride of sodium (salt) 28·05948 ” Chloride of potassium 0·76552 ” Chloride of magnesium 3·66658 ” Bromide of magnesium 0·02929 ” Sulphate of magnesia 2·29578 ” Sulphate of lime 0·40662 ” Carbonate of lime 0·03301 ” (With traces of iodine and ammonia). --—--—---- 1000·00000 grains.
There is much more oxygen in water than in air, as can be ascertained by analysis of these compounds. This great proportion in favour of water enables fish to breathe by passing the water through the gills. Marine animals (not fishes), like the whale,—which is a warm-blooded creature, and therefore not suited to exist without air,—are obliged to come to the surface to breathe. The density of salt water is much greater than that of fresh water, and therefore swimming and flotation is easier in the sea than in a river. We shall have more to say of water by-and-by.
NITROGEN—SYMBOL =N=; ATOMIC WEIGHT 14.
We have already made some reference to this gas when speaking of the atmosphere and its constituents, of which nitrogen is the principal. From its life-destroying properties it is called “azote” by French chemists, and when we wish to obtain a supply of nitrogen all we have to do is to take away the oxygen from the air by burning phosphorus on water under a glass. Nitrogen is not found frequently in solid portions of the globe. It is abundant in animals. It is without colour or smell, and can be breathed in air without danger. It is heavy and sluggish; but if we put a taper into a jar of nitrogen it will go out, and animals die in the gas for want of oxygen, as nitrogen alone cannot support life.
The affinity of nitrogen for other substances is not great, but it gives rise to five compounds, which are as below, in the order they are combined with oxygen:—
Nitrous oxide (“laughing gas”) (Monoxide) N_{2}O. Nitric oxide Dioxide N_{2}O_{2}. Nitrous acid Trioxide N_{2}O_{3}. Nitric peroxide Tetroxide N_{2}O_{4}. Nitric acid Pentoxide N_{2}O_{5}.
These compounds are usually taken as representative examples of combining weight, and as explanatory of the symbolic nomenclature of chemistry, as they advance in such regular proportions of oxygen with nitrogen. The combining weight of nitrogen is 14, and when two parts combine with five of oxygen it makes nitric acid, and we put it down as N_{2}O_{5}; on adding water, HNO_{3}, as we can see by eliminating the constituents and putting in the proportions. Actually it is H_{2}N_{2}O_{6}, or, by division, HNO_{3}.
Nitrogen plays a very important part in nature, particularly in the vegetable kingdom. Nitric acid has been known for centuries. Geber, the alchemist, was acquainted with a substance called “nitric,” which he found would yield a dissolvent under certain circumstances. He called it “dissolving fluid.” At the end of the twelfth century Albert Magnus investigated the properties of this acid, and in 1235 Raymond Lully prepared nitre with clay, and gave the liquid the name of “aqua-fortis.” But till 1849 nitric acid was only known as a hydrate,—that is, in combination with water,—but now we have the anhydrous acid.
Oxygen and nitrogen combine under the influence of electricity, as shown by Cavendish, who passed a current through an atmospheric mixture of oxygen and nitrogen, in a tube terminating in a solution of potash, lime, and soda. Every time the spark passed, the volume of gas diminished, and nitric acid was formed, as it is in thunderstorms, when it does not remain free, but unites with ammonia, and forms a highly useful salt, which promotes vegetable growth. Here is another instance of the usefulness of thunderstorms, and of the grand provisions of nature for our benefit. Nitric acid is obtained by distilling nitre with sulphuric acid. The liquid is, when pure, colourless, and is a powerful oxidizer. It dissolves most metals, and destroys vegetable and animal substances. By an addition of a little sulphuric acid the water is taken from the nitric acid, and a very powerful form of it is the result. The acid is of great use in medicine, and as an application to bites of rabid animals or serpents. It converts cotton waste into “gun-cotton” by a very simple process of steeping, washing, and pressing. From the hydraulic press it comes in discs like “quoits,” which will burn harmlessly and smoulder away, but if detonated they explode with great violence. As a rule, when damp, it is not dangerous, but it can be fired even when wet. It will explode at a less temperature than gunpowder, and, moreover, yields no smoke, nor does it foul a gun. Gun-cotton, when dissolved in ether, gives us collodion for photographic purposes.
In speaking farther of the compounds of nitrogen with oxygen, we will limit ourselves to the monoxide, or laughing gas. This is now used as an anæsthetic in dentistry, etc., and is quite successful, as a rule. People afflicted with heart disease should not use it without advice, however. When inhaled into the lungs it makes the subject very hilarious, and the effect is rather noisy. It is obtained from the nitrate of ammonia, which, on the application of heat, decomposes into nitrous oxide and vapour. Warm water should be used for the trough. The gas is a powerful supporter of combustion.
Binoxide of nitrogen is of importance in the manufacture of sulphuric acid.
Nitrogen combines with hydrogen, forming various compounds. These are the “amines,” also ammonia, and ammonium. Ammonia possesses the properties of a base. Its name is derived from Jupiter Ammon, near whose temple it was prepared, from camels’ dung. But bodies containing nitrogen give off ammonia in course of distilling, and hartshorn is the term applied to horn-cuttings, which yield ammonia, which is a colourless gas of strong odour and taste now obtained from gas-works.
To obtain ammonia heat equal parts of chloride of ammonia (sal ammoniac) and quick-lime powdered (see fig. 357). The gas must be collected over mercury, because it is very soluble in water. Ammonia is useful to restore tipsy people and fainting ladies. A solution of ammonia is used for cauteries. Ammoniacal gas is remarkable for its solubility in water. To prepare the solution the gas is forced through a series of flasks. The tubes carrying the gas should be continued to the bottoms of the flasks, else the solution, being lighter than water, the upper portion alone would be saturated. The tubes carrying away the solution are raised a little, so that the renewal is continually proceeding. The gas liquifies under a pressure of six atmospheres, at a temperature of 10° Cent. This experiment can be artificially performed by heating chloride of silver saturated with ammonia, and the silver will part with the gas at a temperature of 40° C. The gas will then condense in a liquid form in the tube. The experiment may be facilitated by placing the other extremity of the tube in snow and salt, and by the liquid we can obtain intense cold. This experiment has been made use of by M. Carré in his refrigerator (which was described in the Physics’ section), by which he freezes water. We may, however, just refer to the process. Whenever the condition of a body is changed from that of liquid to a gas, the temperature is greatly lowered, because the heat becomes “latent.” The latest freezing machine consists of an apparatus as shown in the illustrations herewith (figs. 359 and 360). The machine is of wrought iron, and contains, when ready for action, a saturated solution of ammonia at zero. This is in communication with another and an air-tight vessel, of which the centre is hollow. The first process is to heat the solution, and the gas escapes into the second “vase,” which is surrounded by cold water, and quite unable to escape. A tremendous pressure is soon obtained, and this, added to the cold water, before long liquifies the ammonia, and when the temperature indicates 130° the hot vessel is suddenly cooled by being put into the water. The gas is thus suddenly converted into a liquid, the water in the second hollow vase is taken out, and the bottle to be frozen is put into the cavity. The cold is so great, in consequence of the transformation of the liquid ammonia into a gas, that it freezes the water in any vessel put into the receiver. The ammonia can be reconverted into liquid and back again, so no loss is occasioned by the process, which is rapid and simple. This is how great blocks of ice are produced in water-bottles.
The one important point upon which care is necessary is the raising of the temperature. If it be elevated beyond 130° C., the pressure will be too great, and an explosion will occur.
The abundant formation of ammonia from decaying animal matter is evident to everyone, and depends upon the presence of moisture to a great extent. Chloride of ammonia is called sal-ammoniac, and the carbonate of ammonia crystallizes from the alkaline liquid produced by the distillation of certain animal matter. The compounds of ammonia are easily recognized by a certain sharp taste. They are highly valuable remedial agents, acting particularly upon the cutaneous system, and when taken internally, produce the effect of powerful sudorifics. Their volatility, and the facility with which they are expelled from other substances, render them of great importance in chemistry, and peculiarly fit them for the purposes of many chemical analyses. The ammonia compounds display a remarkable analogy to the corresponding combinations of potash and soda. The compounds of ammonia are highly important in their relation to the vegetable kingdom. It may be assumed that all the nitrogen of plants is derived from the ammonia which they absorb from the soil, and from the surrounding atmosphere.
The similarity of ammonia to the metallic oxides has led to the conjecture that all its combinations contain a compound metallic body, which has received the name ammonium (NH_{4}); but no one has yet succeeded in its preparation, although by peculiar processes it may be obtained in the form of an amalgam.
Ammonias, in which one or more atoms of hydrogen are replaced by basic radicals, are termed Amides, or Amines.
XXVIII. NON-METALLIC ELEMENTS (continued)
CHLORINE—BROMINE—IODINE—FLUORINE—CARBON—SULPHUR—PHOSPHORUS—SILICON— BORON—TELLURIUM—ARSENIC.
Chlorine (Cl.) is usually found with sodium in the mineral kingdom, and this chloride of sodium is our common salt. Chlorine can be obtained by heating hydrochloric acid with binoxide of manganese. (Atomic weight 35·5.)
Chlorine possesses a greenish-yellow colour, hence its name “Chloros,” green. It should be handled carefully, for it is highly injurious and suffocating. It possesses a great affinity for other substances, and attacks the metals. For hydrogen it has a great affection, and when hydrogen is combined with any other substances chlorine immediately attacks them, and in time destroys them. But even this destructive and apparently objectionable quality makes chlorine very valuable; for if we carry the idea to its conclusion, we shall find that it also destroys offensive and putrid matter, and purifies the atmosphere very much. Most colouring matters include hydrogen, and therefore they are destroyed by chlorine, which is a great “bleacher” as well as a purifier. If we dip any vegetable dyes into a jar of chlorine, they will become white if the dyed substances are damp.
Hydrochloric acid is known as muriatic acid and spirits of salt. It is obtained when salt is treated with sulphuric acid and the gas comes off into water. Equal parts of the acid and the salt are put into a flask as in the cut (fig. 362), and diluted with water. The mixture is then heated. The gas is condensed in the bottles half-full of water. The result gives sulphate of soda and hydrochloric acid. This acid is procured in soda manufactories, and with nitric acid is called “aqua regia,” a solvent for gold. When chlorine and hydrogen are mixed in equal proportions they explode in sunlight. In the dark or by candle-light they are harmless. Dry chlorine gas can be obtained by interposing a glass filled with some chloride of calcium. The gas being heavier than air (about 2½ times), displaces it in the flask, and when it is filled another can be placed in position. This mode causes a little waste of gas, which should not be breathed.
Chlorine possesses a great affinity for certain bodies. If the gas be thrown upon phosphorus, the latter will burn brilliantly. Arsenic, tin, and antimony when powdered and poured from a shoot into a vase of chlorine will burst into brilliant sparks, and other metals will glow when introduced to this gas. Chlorine forms many unstable combinations with oxygen. Its combination with hydrogen has already been referred to.
BROMINE is a rare element. (Symbol Br. Atomic weight 80.)
It is deep brownish red, very volatile, and of a peculiar odour. Bromine unites with the elementary bodies, and forms some oxygen compounds. It resembles chlorine in its properties, and is used in medicine and in photography. It is found in saline springs and in salt water, combined with soda and magnesium. The presence of bromine may easily be detected in the strong smell of seaweed. Its combinations with metals are termed bromides. It is a powerful poison.
IODINE is another relative of chlorine. It is found in seaweed, which by burning is reduced to kelp. When iodine is heated a beautiful violet vapour comes off, and this characteristic has given it its name (“iodes,” violet). Iodine was discovered by Courtois, of Paris, and in 1813, Gay Lussac made it a special study. It is solid at ordinary temperatures, and assumes crystallized forms in plates of metallic lustre. It is an excellent remedy in “goitre” and such affections. (Symbol I. Atomic weight 127.)
FLUORINE is very difficult to prepare. Fluor spar is a compound of fluorine and calcium. This element is gaseous, and combines so rapidly that it is very difficult to obtain in a free state. Etching on glass is accomplished by means of hydrofluoric acid, for fluorine has a great affinity for silicic acid, which is contained in glass. The glass is covered with wax, and the design is traced with a needle. The acid attacks the glass and leaves the wax, so the design is eaten in. (Symbol F. Atomic weight 19.)
Chlorine, fluorine, bromine, and iodine are termed “Halogens” (producers of salts). They appear, as we have seen, in a gaseous, liquid, and solid form respectively.
CARBON is the most, or one of the most, largely diffused elements in nature, and claims more than a passing notice at our hands, though even that must be brief. We may put down carbon next to oxygen as the most important element in the world. The forms assumed by carbon are very variable, and pervade nature in all its phases. We have carbon in crystals, in the animal and vegetable kingdoms, and amongst the chief minerals a solid, odourless, tasteless, infusible, and almost insoluble body. In various combinations carbon meets us at every turn; united with oxygen it forms carbonic acid, which we exhale for the plants to imbibe. We have it in coal, with hydrogen and oxygen. We have it building up animal tissues, and it is never absent in two out of the three great divisions of nature—the plants and the animals (Symbol C; Atomic W. 12).
We have carbon in three different and well-known conditions; as the diamond, as graphite, or black-lead, and as charcoal. The properties of the diamond are well known, and we shall, when we get to Crystallography, learn the forms of diamond or crystals of carbon. At present we give an illustration or two, reserving all explanation for the present. Diamond cutting is a matter of some difficulty, and it requires skill to cut in the proper direction. Diamonds are found in India, Brazil, and at the Cape of Good Hope, in alluvial soil. The identity of diamond and charcoal was discovered accidentally. An experiment to fuse a few small diamonds resulted in their disappearance, and when the residue was examined it was found that the diamonds had been burned, that they had combined with oxygen and formed carbonic acid, just as when coal burns. The diamond is the hardest of all substances, the most valuable of gems, and the purest condition in which carbon appears.
GRAPHITE (Plumbago) is termed “black-lead,” and is the next purest form of carbon. It crystallizes and belongs to the primitive formations. In Cumberland it is dug up and used to make pencils; the operations can be seen at Keswick. It has other uses of a domestic character.
Charcoal is the third form of carbon, and as it possesses no definite form, is said to be amorphous. Charcoal is prepared in air-tight ovens, so that no oxygen can enter and burn the wood thus treated. Coke is the result of the same process applied to coal. The gas manufactories are the chief depôts for this article, and it is used in locomotive engines. The various smokeless coals and prepared fuels, however, are frequently substituted.
Coke ovens were formerly much resorted to by the railway companies, who found the ordinary coal too smoky for locomotive purposes, and apt to give rise to complaints by passengers and residents near the line.
The origin of wood charcoal we have seen. All vegetable substances contain carbon. When we burn wood, in the absence of air as far as possible, oxygen and hydrogen are expelled. The wood is piled in layers as in the illustration (fig. 368), covered over with turf and mould, with occasional apertures for air. This mass is ignited, the oxygen and hydrogen are driven off, and carbon remains. (Animal charcoal is obtained from calcining bones). Wood charcoal attracts vapours, and water, if impure, can be purified by charcoal, and any impure or tainted animal matter can be rendered inoffensive by reason of charcoal absorbing the gases, while the process of decay goes on just the same. Housekeepers should therefore not always decide that meat is good because it is not offensive to the olfactory nerves. Charcoal will remove the aroma, but the meat may be nevertheless bad. The use of charcoal in filters is acknowledged universally, and as a constituent of gunpowder it is important.
Carbon is not easily affected by the atmospheric air, or in the earth; so in many instances wood is charred before being driven into the ground; and casks for water are prepared so. Soot is carbon in a pulverised condition, and Indian ink is manufactured with its assistance.
The preparation of wood charcoal gives occupation to men who are frequently wild and untutored, but the results of their labour are very beneficial. Care should be taken not to sleep in a room with a charcoal stove burning, unless there is ample vent for the carbonic acid gas, for it will cause suffocation. Lampblack is obtained by holding a plate over the flame of some resinous substance, which deposits the black upon it. There is a special apparatus for this purpose.
Carbon combines with oxygen to make carbonic acid gas, as we have already mentioned, and in other proportions to form a more deadly compound than the other. The former is the dioxide (CO_{2}), the latter the monoxide, or carbonic oxide (CO). The dioxide is the more important, being held in the atmosphere, and combined with lime in chalk. All sparkling beverages contain carbonic acid, to which their effervescence is due. The soda and other mineral waters owe their sparkle to this gas. Soda-water consists of a weak solution of carbonate of soda and the acid. There is a vessel holding chalk and water, and another containing some sulphuric acid. When the sulphuric acid is permitted to unite with the chalk and water, carbonic acid is liberated. A boy turning a wheel forces the gas into the water in the bottles, or the water and carbonate of soda is drawn off thus impregnated into bottles and corked down, in the manner so familiar to all. The bottles are made of the shape depicted, so that the bubble of air shall be at the top when the bottle lies down. If it be not kept so, the air will eventually escape, no matter how tightly the cork be put in. The ordinary “soda-water” contains scarcely any soda. It is merely water, chalk, and carbonic acid. The “Gazogene” is made useful for small quantities of soda-water, and is arranged in the following manner. The appearance of it is familiar to all. It consists of a double vessel, into the upper part of which a solution of any kind—wine and water, or even plain water—is put, to be saturated with carbonic acid, or “aerated,” and into the lower one some carbonate of soda and tartaric acid. A tube leads from this lower to the top of the upper vessel, which screws on and off. By shaking the apparatus when thus charged and screwed together, some of the liquid descends through the tube into the lower vessel and moistens the soda and acid, which therefore act on each other, and cause carbonic acid to be disengaged; this, rising up through the tube (which is perforated with small holes at the upper part), disperses itself through the liquid in small bubbles, and causes sufficient pressure to enable the liquid to absorb it, which therefore effervesces when drawn off by the tap.
Carbonic acid can be liquified, and then it is colourless. In a solid form it resembles snow, and if pressed with the fingers it will blister them. Being very heavy the gas can be poured into a vase, and if there be a light in the receptacle the flame will be immediately extinguished.
That even the gas introduced into seltzer-water is capable of destroying life, the following experiment will prove. Let us place a bird within a glass case as in the illustration (fig. 373), and connect the glass with a bottle of seltzer-water or a siphon. As soon as the liquid enters, the carbonic acid will ascend, and this, if continued for a long time, would suffocate the bird, which soon begins to develop an appearance of restlessness.
We have already remarked upon the important part taken by this gas in nature, so we need only mention its existence in pits and caves. There are many places in which the vapour is so strong as to render the localities uninhabitable. In the Middle Ages the vapours were attributed to the presence of evil spirits, who were supposed to extinguish miners’ lamps, and suffocate people who ventured into the caves. In the Grotto Del Cane there is still an example, and certain caves of Montrouge are often filled with the gas. A lighted taper held in the hand will, by its extinction, give the necessary warning. Oxygen and carbon are condensed in carbonic acid, for the gas contains a volume of oxygen equal to its own. If we fill a glass globe, as per illustration (fig. 374), with pure oxygen, and in the globe insert two carbon points, through which we pass a current of electricity, we shall find, after the experiment, that if the stop-cock be opened, there is no escape of gas, and yet the mercury does not rise in the tube, so the oxygen absorbed has been replaced by an equal volume of carbonic acid.
The other combination of carbon with oxygen is the carbonic oxide (CO), and when a small quantity of oxygen is burnt with it it gives a blue flame, as on the top of the fire in our ordinary grates. This gas is present in lime kilns, and is a very deadly one. We must now pass rapidly through the compounds of carbon with hydrogen, merely referring to coal for a moment as we go on.
Coal, of which we shall learn more in Mineralogy and Geology, is a combination, mechanical or otherwise, and is the result of the decomposition of vegetable matter in remote ages,—the so-called “forests,” which were more like the jungles than the woods of the present day. Moss and fern played prominent parts in this great transformation, as we can see in the Irish peat-bogs, where the first steps to the coal measures are taken.
The compounds of carbon with hydrogen are important. There is the “light” carburetted hydrogen (CH_{4}), which is usually known as fire-damp in coal mines. It is highly inflammable and dangerous. The safety-lamp invented by Davy is a great protection against it, for as the gas enters it is cooled by the wire, and burns within harmlessly. The explosion warns the miner. “Heavy” carburetted hydrogen possesses double the quantity of carbon (C_{2}H_{4}). It is also explosive when mixed with oxygen.
The most useful compound is coal-gas, and though its principal function appears to be in some manner superseded by electricity, “gas” is still too important to be put aside. It can easily be obtained by putting small fragments of coal in the bowl of a tobacco-pipe, closing the bowl with clay, and putting it in the fire. Before long the gas will issue from the stem of the pipe, and may either be lighted or collected in a bladder. For the use of the “million,” however, gas is prepared upon a very large scale, and is divided into three processes—its “formation,” “purification,” and its “collection” for distribution to consumers. The first process is carried on by means of retorts shown in the illustration (fig. 376). The first portion of the next figure is a section of a furnace, the other part shows two furnaces from the front. The following is the mode employed. The coal is put into retorts fitted to the furnace, so that they are surrounded by the flames, and terminating in a horizontal tube called the hydraulic main, E, which is in its turn connected with a pit or opening for the reception of the tar and ammoniacal liquor, etc., which condenses from the gas. It then passes up and down a series of tubes in water, called a “condenser,” and in this are reservoirs or receptacles for any tar and ammonia that remain. But sulphur is still present, so the gas is carried to the purifying apparatus (D in fig. 378), which consists of a large cylindrical vessel air-tight, with an inverted funnel, nearly filled with a mixture of lime and water. The gas bubbles in, and the sulphur unites with the lime, while the gas rises to the top (trays of lime are used when the gas enters from the bottom). The Gasometer, a large vessel closed at the top and open below, dips into a large trough of circular shape. The gasometer is balanced by weights and chains, and may be raised (See fig. 379). When quite empty the top rests upon the ground, and when the gas enters it is raised to the top of the frame which supports it. We have now our Gasometer full. When the time comes to fill the pipes for lighting purposes, some of the weights are removed, the Gasometer falls down slowly, and forces the gas through the tubes into the main supply to be distributed. About four cubic feet of gas is obtained from every pound of coal. When gas and air become mixed, the mixture is very explosive. In a house where an escape of gas is detected let the windows be opened at the top, and no light introduced for several minutes.
It has been calculated that one ton of good coal produces the following:—
1 Chaldron of coke weighing 1,494 lbs. 12 Gallons of tar ” 135 lbs. 12 Gallons of ammoniacal liquor ” 100 lbs. 5,900 Cubic feet of gas ” 291 lbs. Loss (water) ” 220 lbs. ---------- Total 2,240 lbs. ==========
We can thus estimate the profits of our gas companies at leisure. The analysis of gas made by Professor Bunsen is as under, in 100 parts.
Hydrogen 45·58 Marsh gas 34·90 Carbonic oxide 6·64 Olefiant gas 4·08 Butyline 2·38 Sulphide of hydrogen 0·29 Nitrogen 2·46 Carbonic acid 3·67 ------ 100·00 ======
Gas, therefore, is very injurious, for it rapidly vitiates the atmosphere it burns in, and is very trying to the eyes, as well as destructive to gilt ornaments.
Tar is familiar to all readers, and though unpleasant to handle or to smell, it produces the beautiful aniline dyes. Tar pills are very efficacious for some blood disorders, and will remove pimples, etc., from the face, and cure “boils” effectually. If a dose of five be taken first, in a day or two four, and so on, no second remedy need be applied. We have known cases finally cured, and no recurrence of boils ever ensued after this simple remedy.
Tar is one of the results left in the distillation both of wood and coal: in places where wood is plentiful and tar in request, it is produced by burning the wood for that purpose; and in some of the pits in which charcoal is produced, an arrangement is made to collect the tar also. Coal-tar and wood-tar are different in some respects, and are both distilled to procure the napthas which bear their respective names. From wood-tar creosote is also extracted, and it is this substance which gives the peculiar tarry flavours to provisions, such as ham, bacon, or herrings, cured or preserved by being smoked over wood fires. Tar is used as a sort of paint for covering wood-work and cordage when much exposed to wet, which it resists better than anything else at the same price; but the tar chiefly used for these purposes is that produced by burning fir or deal wood and condensing the tar in a pit below the stack of wood; it is called Stockholm tar, as it comes chiefly from that place.
Carbon only combines with nitrogen under peculiar circumstances. This indirect combination is termed cyanogen (CN). It was discovered by Gay-Lussac, and is used for the production of Prussian blue. Hydrocyanide of potassium (Prussic acid) is prepared by heating cyanide of potassium with sulphuric acid. It is a deadly poison, and found in peach-stones. Free cyanogen is a gas. The bisulphide of carbon is a colourless, transparent liquid. It will easily dissolve sulphur and phosphorus and several resins. When phosphorus is dissolved in it, it makes a very dangerous “fire,” and one difficult to extinguish. We must now leave carbon and its combinations, and come to sulphur.
SULPHUR is found in a native state in Sicily and many other localities which are volcanic. It is a yellow, solid body, and as it is never perfectly free from earthy matter, it must be purified before it can be used. It possesses neither taste nor smell, and is insoluble in water. Sulphur is purified in a retort, C D, which communicates with a brick chamber, A. The retort is placed over a furnace, K, and the vapour passes into the chimney through the tube, D, where it condenses into fine powder called “flowers of sulphur” (brimstone). A valve permits the heated air to pass off, while no exterior air can pass in, for explosions would take place were the heated vapour to meet the atmospheric air. The danger is avoided by putting an air reservoir outside the chimney which is heated by the furnace. The sulphur is drawn out through the aperture, r, when deposited on the floor of the chamber. The sulphur is cast into cylinders and sold. Sulphur is soluble in bisulphide of carbon, and is used as a medical agent.
The compounds of sulphurs with oxygen form an interesting series. There are two anhydrous oxides (anhydrides),—viz., sulphurous and sulphuric anhydride (SO_{2} and SO_{3}). There are two notable acids formed by the combination with water, sulphurous and sulphuric, and some others, which, as in the case of nitrogen, form a series of multiple proportions, the oxygen being present in an increasing regularity of progression, as follows:—
Name of Acid. Chemical Formula.
Hypo-sulphurous acid H_{2}SO_{2} Sulphurous acid H_{2}SO_{3} Sulphuric acid H_{2}SO_{4} Thio-sulphuric, or hypo-sulphuric acid H_{2}S_{2}O_{3} Dithionic acid H_{2}S_{2}O_{6} Trithionic acid H_{2}S_{3}O_{6} Tetrathionic acid H_{2}S_{4}O_{6} Pentathionic acid H_{2}S_{5}O_{6}
The last four are termed “polythionic,” because the proportions of sulphur vary with constant proportions of the other constituents.
The sulphurous anhydride mentioned above is produced when we burn sulphur in the air, or in oxygen; it may be obtained in other ways. It is a colourless gas, and when subjected to pressure may be liquified, and crystallized at very low temperature. It was formerly called sulphuric acid. It is a powerful “reducing agent,” and a good antiseptic. It dissolves in water, and forms the H_{2}SO_{3}, now known as sulphurous acid.
Sulphuric acid is a most dangerous agent in wicked or inexperienced hands, and amateurs should be very careful when dealing with it. It takes the water from the moist air, and from vegetable and animal substances. It carbonizes and destroys all animal tissues. Its discovery is due to Basil Valentine, in 1440. He distilled sulphate of iron, or green vitriol, and the result was “oil of vitriol.” It is still manufactured in this way in the Hartz district, and the acid passes by retorts into receivers. The earthen retorts, A, are arranged in the furnace as in the illustration, and the receivers, B, containing a little sulphuric acid, are firmly fixed to them. The oily brown product fumes in the air, and is called “fuming sulphuric acid,” or Nordhausen acid. Sulphuric acid is very much used in chemical manufactures, and the prices of many necessaries, such as soap, soda, calico, stearin, paper, etc., are in close relationship with the cost and production of sulphur, which also plays an important part in the making of gunpowder. The manufacture of the acid is carried on in platinum stills.
Sulphuretted hydrogen, or the hydric sulphide (H_{2}S), is a colourless and horribly-smelling gas, and arises from putrefying vegetable and animal matter which contains sulphur. The odour of rotten eggs is due to this gas, which is very dangerous when breathed in a pure state in drains, etc. It can be made by treating a sulphide with sulphuric acid. It is capable of precipitating the metals when in solution, and so by its aid we can discover the metallic ingredient if it be present. The gas is soluble in water, and makes its presence known in certain sulphur springs. The colour imparted to egg-spoons and fish-knives and forks sometimes is due to the presence of metallic sulphides. The solution is called hydro-sulphuric acid.
PHOSPHORUS occurs in very small quantities, though in the form of phosphates we are acquainted with it pretty generally, and as such it is absorbed by plants, and is useful in agricultural operations. In our organization—in the brain, the nerves, flesh, and particularly in bones—phosphorus is present, and likewise in all animals. Nevertheless it is highly poisonous. It is usually obtained from the calcined bones of mammalia by obtaining phosphoric acid by means of acting upon the bone-ash with sulphuric acid. Phosphorus when pure is colourless, nearly transparent, soft, and easily cut. It has a strong affinity for oxygen. It evolves white vapour in atmospheric air, and is luminous; to this element is attributable the luminosity of bones of decaying animal matter. It should be kept in water, and handled—or indeed not handled—but grasped with a proper instrument.
Phosphorus is much used in the manufacture of lucifer matches, and we are all aware of the ghastly appearance and ghostly presentment it gives when rubbed upon the face and hands in the dark. In the ripples of the waves and under the counter of ships at sea, the phosphorescence of the ocean is very marked. In Calais harbour we have frequently noticed it of a very brilliant appearance as the mail steamer slowly came to her moorings. This appearance is due to the presence of phosphorus in the tiny animalculæ of the sea. It is also observable in the female glow-worm, and the “fire-fly.” Phosphorus was discovered by Brandt in 1669.
It forms two compounds with oxygen-phosphorous acid, H_{2}PO_{4}, and phosphoric acid, H_{3}PO_{4}. The compound with hydrogen is well marked as phosphuretted hydrogen, and is a product of animal and vegetable decomposition. It may frequently be observed in stagnant pools, for when emitted it becomes luminous by contact with atmospheric air. There is a very pretty but not altogether safe experiment to be performed when phosphuretted hydrogen has been prepared in the following manner. Heat small pieces of phosphorus with milk of lime or a solution of caustic potash; or make a paste of quick-lime and phosphorus, and put into the flask with some quick-lime powdered. Fix a tube to the neck, and let the other end be inserted in a basin of water. (See illustration, fig. 388.) Apply heat; the phosphuretted hydrogen will be given off, and will emerge from the water in the basin in luminous rings of a very beautiful appearance. The greatest care should be taken in the performance of this very simple experiment. No water must on any account come in contact with the mixture in the flask. If even a drop or two find its way in through the bent tube a tremendous explosion will result, and then the fire generated will surely prove disastrous. The experiment can be performed in a cheaper and less dangerous fashion by dropping phosphate of lime into the basin. We strongly recommend the latter course to the student unless he has had some practice in the handling of these inflammable substances, and learnt caution by experience. The effect when the experiment is properly performed is very good, the smoke rising in a succession of coloured rings.
SILICON is not found in a free state in nature, but, combined with oxygen, as Silica it constitutes the major portion of our earth, and even occurs in wheat stalks and bones of animals. As flint or quartz (see Mineralogy) it is very plentiful, and in its purest form is known as rock crystal, and approaches the form of carbon known as diamond. When separated from oxygen, silicon is a powder of greyish-brown appearance, and when heated in an atmosphere of oxygen forms silicic “acid” again, which, however, is not acid to the taste, and is also termed “silica,” or “silex.” It is fused with great difficulty, but enters into the manufacture of glass in the form of sand. The chemical composition of glass is mixed silicate of potassium or sodium, with silicates of calcium, lead, etc. Ordinary window-glass is a mixture of silicates of sodium and calcium; crown glass contains calcium and silicate of potassium. Crystal glass is a mixture of the same silicate and lead. Flint glass is of a heavier composition. Glass can be coloured by copper to a gold tinge, blue by cobalt, green by chromium, etc. Glass made on a large scale is composed of the following materials, according to the kind of glass that is required.
Flint glass (“crystal”) is very heavy and moderately soft, very white and bright. It is essentially a table-glass, and was used in the construction of the Crystal Palace. Its composition is—pure white sea-sand, 52 parts, potash 14 parts, oxide of lead, 34 parts = 100.
Plate Glass. Crown Glass. Green (Bottle) Glass.
Pure white sand 55 parts. Fine sand 63 parts. Sea sand 80 parts. Soda 35 ” Chalk 7 ” Salt 10 ” Nitre 8 ” Soda 30 ” Lime 10 ” Lime 2 ” --- --- --- 100 ” 100 ” 100 ”
The ingredients to be made into glass (of whatever kind it may be) are thoroughly mixed together and thrown from time to time into large crucibles placed in a circle, A A (fig. 389), in a furnace resting on buttresses, B B, and heated to whiteness by means of a fire in the centre, C, blown by a blowing machine, the tube of which is seen at D. This furnace is shown in prospective in fig. 390. The ingredients melt and sink down into a clear fluid, throwing up a scum, which is removed. This clear glass in the fused state is kept at a white heat till all air-bubbles have disappeared; the heat is then lowered to a bright redness, when the glass assumes a consistence and ductility suitable to the purposes of the “blower.”
Glass blowing requires great care and dexterity, and is done by twirling a hollow rod of iron on one end of which is a globe of melted glass, the workman blowing into the other end all the time. By reheating and twirling a sheet of glass is produced. Plate glass is formed by pouring the molten glass upon a table with raised edges. When cold it is ground with emery powder, and then polished by machinery.
Many glass articles are cast, or “struck-up,” by compression in moulds, and are made to resemble cut-glass, but they are much inferior in appearance. The best are first blown, and afterwards cut and polished. Of whatever kind of glass the article may be, it is so brittle that the slightest blow would break it, a bad quality which is got rid of by a process called “annealing,” that is, placing it while quite hot on the floor of an oven, which is allowed to cool very gradually. This slow cooling takes off the brittleness, consequently articles of glass well annealed are very much tougher than others, and will scarcely break in boiling water.
The kind generally used for ornamental cutting is flint-glass. Decanters and wine-glasses are therefore made of it; it is very bright, white, and easily cut. The cutting is performed by means of wheels of different sizes and materials, turned by a treadle, as in a common lathe, or by steam power; some wheels are made of fine sandstone, some of iron, others of tin or copper; the edges of some are square,or round, or sharp. They are used with sand and water, or emery and water, stone wheels with water only.
In a soluble form silicic acid is found in springs, and thus enters into the composition of most plants and grasses, while the shells and scales of “infusoria” consist of silica. As silicate of alumina,—i.e., clay,—it plays a very important rôle in our porcelain and pottery works.
* * * * *
BORON is found in volcanic districts, in lakes as boracic acid, in combination with oxygen. It is a brownish-green, insoluble powder, in a free state, but as boracic acid it is white. It is used to colour fireworks with the beautiful green tints we see. Soda and boracic acid combine to make borax (or biborate of soda). Another and inferior quality of this combination is tinkal, found in Thibet. Borax is much used in art and manufactures, and in glazing porcelain. (Symbol B, Atomic Weight 11).
SELENIUM is a very rare element. It was found by Berzilius in a sulphuric-acid factory. It is not found in a free state in nature. It closely resembles sulphur in its properties. Its union with hydrogen produces a gas, seleniuretted hydrogen, which is even more offensive than sulphuretted hydrogen. (Symbol Se, Atomic Weight 79).
TELLURIUM is also a rare substance generally found in combination with gold and silver. It is like bismuth, and is lustrous in appearance. Telluretted hydrogen is horrible as a gas. Tellurium, like selenium, sulphur, and oxygen, combines with two atoms of hydrogen. (Symbol Te, Atomic Weight 129).
ARSENIC, like tellurium, possesses many attributes of a metal, and on the other hand has some resemblance to phosphorus. Arsenic is sometimes found free, but usually combined with metals, and is reduced from the ores by roasting; and uniting with oxygen in the air, is known as “white arsenic.” The brilliant greens on papers, etc., contain arsenic, and are poisonous on that account. Arsenic and hydrogen unite (as do sulphur and hydrogen, etc.), and produce a foetid gas of a most deadly quality. This element also unites with sulphur. If poured into a glass containing chlorine it will sparkle and scintillate as in the illustration (fig. 395). (Symbol As, Atomic Weight 75).
Before closing this division, and passing on to a brief review of the METALS, we would call attention to a few facts connected with the metalloids we have been considering. Some, we have seen, unite with hydrogen only, as chlorine; some with two atoms of hydrogen, as oxygen, sulphur, etc., and some with three, as nitrogen and phosphorus; some again with four, as carbon and silicon. It has been impossible in the pages we have been able to devote to the Metalloids to do more than mention each briefly and incompletely, but the student will find sufficient, we trust, to interest him, and to induce him to search farther, while the general reader will have gathered some few facts to add to his store of interesting knowledge. We now pass on to the Metals.
XXIX. THE METALS
WHAT METALS ARE—CHARACTERISTICS AND GENERAL PROPERTIES OF METALS—CLASSIFICATION—SPECIFIC GRAVITY—DESCRIPTIONS.
We have learnt that the elements are divided into metalloids and metals, but the line of demarcation is very faint. It is very difficult to define what a metal is, though we can say what it is not. It is indeed impossible to give any absolute definition of a metal, except as “an element which does not unite with hydrogen, or with another metal to form a chemical compound.” This definition has been lately given by Mr. Spencer, and we may accept it as the nearest affirmative definition of a metal, though obviously not quite accurate.
A metal is usually supposed to be solid, heavy, opaque, ductile, malleable, and tenacious; to possess good conducting powers for heat and electricity, and to exhibit a certain shiny appearance known as “metallic lustre.” These are all the conditions, but they are by no means necessary, for very few metals possess them all, and many non-metallic elements possess several. The “alkali” metals are lighter than water; mercury is a fluid. The opacity of a mass is only in relation to its thickness, for Faraday beat out metals into plates so thin that they became transparent. All metals are not malleable, nor are they ductile. Tin and lead, for example, have very little ductility or tenacity, while bismuth and antimony have none at all. Carbon is a much better conductor of electricity than many metals in which such power is extremely varied. Lustre, again, though possessed by metals, is a characteristic of some non-metals. So we see that while we can easily say what is not a metal, we can scarcely define an actual metal, nor depend upon unvarying properties to guide us in our determination.
The affinity of metals for oxygen is in an inverse ratio to their specific gravity, as can be ascertained by experiment, when the heaviest metal will be the least ready to oxidise. Metals differ in other respects, and thus classification and division become easier. The fusibility of metals is of a very wide range, rising from a temperature below zero to the highest heat obtainable in the blow-pipe, and even then in the case of osmium there is a difficulty. While there can be no question that certain elements, iron, copper, gold, silver, etc., are metals proper, there are many which border upon the line of demarcation very closely, and as in the case of arsenic even occupy the debatable land.
SPECIFIC GRAVITY is the relation which the weight of substance bears to the weight of an equal volume of water, as already pointed out in PHYSICS. The specific gravities of the metals vary very much, as will be seen from the table following—water being, as usual, taken as 1:—
Aluminium 2·56 Antimony 6·7 Arsenic 6· Bismuth 9·7 Cadmium 8·6 Calcium 1·5 Chromium 6·8 Cobalt 8·9 Copper 8·9 Gold 19·3 Indium 7·3 Iridium 21·1 Iron 7·8 Lead 11·3 Lithium ·593 Magnesium 1·74 Manganese 8· Mercury 13·5 Molybdenum 8·6 Nickel 8·8 Osmium 21·4 Palladium 11·8 Platinum 21·5 Potassium ·865 Rhodium 12·1 Rubidium 1·5 Ruthenium 11·4 Silver 10·5 Sodium ·972 Strontium 2·5 Thallium 11·8 Tin 7·2 Titanium 5·3 Tungsten 17·6 Uranium 18·4 Zinc 7·1 Zircon 4·3
Some metals are therefore lighter and some heavier than water.
The table underneath gives the approximate fusing points of some of the metals (Centigrade Scale)—
(Ice melts at 0°.)
Platinum[21] about 1500° Gold ” 1200° Silver ” 1000° Cast iron 1000-1200° Wrought iron ” 1500° Copper ” 1100° Antimony ” 432° Zinc ” 400° Lead ” 330° Bismuth ” 265° Tin ” 235° Sodium ” 97° Potassium ” 60° Mercury ” 40°
There are some metals which, instead of fusing,—that is, passing from the solid to the liquid state,—go away in vapour. These are volatile metals. Mercury, potassium, and sodium, can be thus distilled. Some do not expand with heat, but contract (like ice), antimony and bismuth, for instance, while air pressure has a considerable effect upon the fusing point. Some vaporise at once without liquefying; others, such as iron, become soft before melting.
ALLOYS are combinations of metals which are used for many purposes, and become harder in union. Amalgams are alloys in which mercury is one constituent. Some of the most useful alloys are under-stated:—
Name of Alloy. Composition.
Aluminium bronze Copper and aluminium. Bell metal Copper and tin. Bronze ” ” Gun metal ” ” Brass Copper and zinc. Dutch metal ” ” Mosaic gold ” ” Ormulu ” ” Tombac ” ” German silver Copper, nickel, and zinc. Britannia metal Antimony and tin. Solder ” ” Pewter Tin and lead. Type metal Lead and antimony (also copper at times). Shot Lead and arsenic. Gold currency Gold and copper. Silver currency Silver and copper. Stereotype metal Lead, antimony, and bismuth. Metals combine with chlorine, and produce chlorides, Metals ” ” sulphur ” ” sulphides, Metals ” ” oxygen ” ” oxides, and so on.
The metals may be classed as follows in divisions:—
Metals of the alkalies as POTASSIUM, SODIUM, LITHIUM, AMMONIUM. Metals of the alkaline earths as BARIUM, CALCIUM, MAGNESIUM, STRONTIUM. Metals of the earths as {ALUMINIUM, Cerium, Didymium, Erbium, {Glucinium, Lanthanum, Terbium, {Thorium, Yttrium, Zirconium. Metals proper— Common Metals as {IRON, MANGANESE, COBALT, NICKEL, {COPPER, Bismuth, LEAD, TIN, ZINC, {CHROMIUM, Antimony. Noble Metals as {MERCURY, SILVER, GOLD, PLATINUM, {Palladium, Rhodium, Ruthenium, {Osmium, Iridium.
We cannot attempt an elaborate description of all the metals, but we will endeavour to give a few particulars concerning the important ones, leaving many parts for Mineralogy to supplement and enlarge upon. We shall therefore mention only the most useful of the metals in this place. We will commence with POTASSIUM.
METALS OF THE ALKALIES.
POTASSIUM has a bright, almost silvery, appearance, and is so greatly attracted by oxygen that it cannot be kept anywhere if that element be present—not even in water, for combustion will immediately ensue on water; and in air it is rapidly tarnished. It burns with a beautiful violet colour, and a very pretty experiment may easily be performed by throwing a piece upon a basin of water. The fragment combines with the oxygen of the water, the hydrogen is evolved, and burns, and the potassium vapour gives the gas its purple or violet colour. The metal can be procured by pulverizing carbonate of potassium and charcoal, and heating them in an iron retort. The vapour condenses into globules in the receiver, which is surrounded by ice in a wire basket. It must be collected and kept in naphtha, or it would be oxidised. Potassium was first obtained by Sir Humphrey Davy in 1807. Potash is the oxide of potassium, and comes from the “ashes” of wood.
The compounds of potassium are numerous, and exist in nature, and by burning plants we can obtain potash (“pearlash”). Nitrate of potassium, or nitre (saltpetre), (KNO_{3}), is a very important salt. It is found in the East Indies. It is a constituent of gunpowder, which consists of seventy-five parts of nitre, fifteen of charcoal, and ten of sulphur. The hydrated oxide of potassium, or “caustic potash” (obtained from the carbonate), is much used in soap manufactories. It is called “caustic” from its property of cauterizing the tissues. Iodide, bromide, and cyanide of potassium, are used in medicine and photography.
Soap is made by combining soda (for hard soap), or potash (for soft soap), with oil or tallow. Yellow soap has turpentine, and occasionally palm oil, added. Oils and fats combine with metallic oxides, and oxide of lead with olive oil and resin forms the adhesive plaister with which we are all familiar when the mixture is spread upon linen. Fats boiled with potash or soda make soaps; the glycerine is sometimes set free and purified as we have it. Sometimes it is retained for glycerine soap. Fancy soap is only common soap coloured. White and brown Windsor are the same soap—in the latter case browned to imitate age! Soap is quite soluble in spirits, but in ordinary water it is not so greatly soluble, and produces a lather, owing to the lime in the water being present in more or less quantity, to make the water more or less “hard.”
SODIUM is not unlike potassium, not only in appearance, but in its attributes; it can be obtained from the carbonate, as potassium is obtained from its carbonate. Soda is the oxide of sodium, but the most common and useful compound of sodium is the chloride, or common salt, which is found in mines in England, Poland, and elsewhere. Salt may also be obtained by the evaporation of sea water. Rock salt is got at Salzburg, and the German salt mines and works produce a large quantity. The Carbonate of Soda is manufactured from the chloride of sodium, although it can be procured from the salsoda plants by burning. The chloride of sodium is converted into sulphate, and then ignited with carbonate of lime and charcoal. The soluble carbonate is extracted in warm water, and sold in crystals as soda, or (anhydrous) “soda ash.” The large quantity of hydrochloric acid produced in the first part of the process is used in the process of making chloride of lime. A few years back, soda was got from Hungary and various other countries where it exists as a natural efflorescence on the shores of some lakes, also by burning sea-weeds, especially the common bladder wrack (Fucus vesiculosus), the ashes of which were melted into masses, and came to market in various states of purity. The bi-carbonate of soda is obtained by passing carbonic acid gas over the carbonate crystals. Soda does not attract moisture from the air. It is used in washing, in glass manufactories, in dyeing, soap-making, etc.
Sulphate of Soda is “Glauber’s Salt”; it is also employed in glass-making. Mixed with sulphuric acid and water, it forms a freezing mixture. Glass, as we have seen, is made with silicic acid (sand), soda, potassa, oxide of lead, and lime, and is an artificial silicate of soda.
LITHIUM is the lightest of metals, and forms the link between alkaline and the alkaline earth metals. The salts are found in many places in solution. The chloride when decomposed by electricity yields the metal.
CÆSIUM and RUBIDIUM require no detailed notice from us. They were first found in the solar spectrum, and resemble potassium.
AMMONIUM is only a conjectural metal. Ammonia, of which we have already treated, is so like a metallic oxide that chemists have come to the conclusion that its compounds contain a metallic body, which they have named hypothetically AMMONIUM. It is usually classed amongst the alkaline metals. The salts of ammonia are important, and have already been mentioned. Muriate (chloride) of ammonia, or sal-ammoniac, is analogous to chloride of sodium and chloride of potassium. It is decomposed by heating it with slaked lime, and then gaseous ammonia is given off.
THE METALS OF THE ALKALINE EARTHS.
BARIUM is the first of the four metals we have to notice in this group, and will not detain us long, for it is little known in a free condition. Its most important compound is heavy spar (sulphate of baryta), which, when powdered, is employed as a white paint. The oxide of barium, BaO, is termed baryta.
Nitrate of Baryta is used for “green fire,” which is made as follows:—Sulphur, twenty parts; chlorate of potassium, thirty-three parts; and nitrate of baryta, eighty parts (by weight).
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CALCIUM forms a considerable quantity of our earth’s crust. It is the metal of lime, which is the oxide of calcium. In a metallic state it possesses no great interest, but its combinations are very important to us. Lime is, of course, familiar to all. It is obtained by evolving the carbonic acid from carbonate of lime (CaO).
The properties of this lime are its white appearance, and it develops a considerable amount of heat when mixed with water, combining to make hydrate of lime, or “slaked lime.” This soon crumbles into powder, and as a mortar attracts the carbonic acid from the air, by which means it assumes the carbonate and very solid form, which renders it valuable for cement and mortar, which, when mixed with sand, hardens. Caustic lime is used in whitewashing, etc.
Carbonate of Lime (CaCO_{3}) occurs in nature in various forms, as limestone, chalk, marble, etc. Calc-spar (arragonite) is colourless, and occurs as crystals. Marble is white (sometimes coloured by metallic oxides), hard, and granular. Chalk is soft and pulverizing. It occurs in mountainous masses, and in the tiniest shells, for carbonate of lime is the main component of the shells of the crustacea, of corals, and of the shell of the egg; it enters likewise into the composition of bones, and hence we must regard it as one of the necessary constituents of the food of animals. It is an almost invariable constituent of the waters we meet with in Nature, containing, as they always do, a portion of carbonic acid, which has the power of dissolving carbonate of lime. But when gently warmed, the volatile gas is expelled, and the carbonate of lime deposited in the form of white incrustations upon the bottom of the vessel, which are particularly observed on the bottoms of tea-kettles, and if the water contains a large quantity of calcareous matter, even our water-bottles and drinking-glasses become covered with a thin film of carbonate of lime. These depositions may readily be removed by pouring into the vessels a little dilute hydrochloric acid, or some strong vinegar, which in a short time dissolves the carbonate of lime.
Sulphate of Lime (CaSO_{4}) is found in considerable masses, and is commonly known under the name of Gypsum. It occurs either crystallized or granulated, and is of dazzling whiteness; in the latter form it is termed Alabaster, which is so soft as to admit of being cut with a chisel, and is admirably adapted for various kinds of works of art. Gypsum contains water of crystallization, which is expelled at a gentle heat. But when ignited, ground, and mixed into a paste with water, it acquires the property of entering into chemical combination with it, and forming the original hydrate, which in a short time becomes perfectly solid. Thus it offers to the artist a highly valuable material for preparing the well-known plaster of Paris figures, and by its use the noblest statues of ancient and modern art have now been placed within the reach of all. Gypsum, moreover, has received a valuable application as manure. In water it is slightly soluble, and imparts to it a disagreeable and somewhat bitterish, earthy taste. It is called “selenite” when transparent.
Phosphate of Lime constitutes the principal mass of the bones of animals, and is extensively employed in the preparation of phosphorus; in the form of ground bones it is likewise used as a manure. It appears to belong to those mineral constituents which are essential to the nutrition of animals. It is found in corn and cereals, and used in making bread; so we derive the phosphorus which is so useful to our system.
Chloride of Lime is a white powder smelling of chlorine, and is produced by passing the gas over the hydrate of lime spread on trays for the purpose. It is the well-known “bleaching powder.” It is also used as a disinfectant. The Fluoride of Calcium is Derbyshire spar, or “Blue John.” Fluor spar is generally of a purple hue. We may add that hard water can be softened by adding a little powdered lime to it.
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MAGNESIUM sometimes finds a place with the other metals, for it bears a resemblance to zinc. Magnesium may be prepared by heating its chloride with sodium. Salt is formed, and the metal is procured. It burns very brightly, and forms an oxide of magnesia (MgO). Magnesium appears in the formation of mountains occasionally. It is ductile and malleable, and may be easily melted.
Carbonate of Magnesia, combining with carbonate of lime, form the Dolomite Hills. When pure, the carbonate is a light powder, and when the carbonic acid is taken from it by burning it is called Calcined Magnesia.
The Sulphate of Magnesia occurs in sea-water, and in saline springs such as Epsom. It is called “Epsom Salts.” Magnesium wire burns brightly, and may be used as an illuminating agent for final scenes in private theatricals. Magnesite will be mentioned among Minerals.
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STRONTIUM is a rare metal, and is particularly useful in the composition of “red-fire.” There are the carbonate and sulphate of strontium; the latter is known as Celestine. The red fire above referred to can be made as follows, in a dry mixture. Ten parts nitrate of strontia, 1½ parts chlorate of potassium, 3½ parts of sulphur, 1 part sulphide of antimony, and ½ part charcoal. Mix well without moisture, enclose in touch paper, and burn. A gorgeous crimson fire will result.
METALS OF THE EARTHS.
ALUMINIUM (Aluminum) is like gold in appearance when in alloy with copper, and can be procured from its chloride by decomposition with electricity. It occurs largely in nature in composition with clays and slates. Its oxide, alumina (Al_{2}O_{3}), composes a number of minerals, and accordingly forms a great mass of the earth. Alumina is present in various forms (see Minerals) in the earth, all of which will be mentioned under Crystallography and Mineralogy. The other nine metals in this class do not call for special notice.
HEAVY METALS
IRON, which is the most valuable of all our metals, may fitly head our list. So many useful articles are made of it, that without consideration any one can name twenty. The arts of peace and the glories of war are all produced with the assistance of iron, and its occurrence with coal has rendered us the greatest service, and placed us at the head of nations. It occurs native in meteoric stones.
Iron is obtained from certain ores in England and Sweden, and these contain oxygen and iron (see Mineralogy). We have thus to drive away the former to obtain the latter. This is done by putting the ores in small pieces into a blast furnace (fig. 402) mixed with limestone and coal. The process of severing the metal from its ores is termed smelting, the air supplied to the furnace being warmed, and termed the “hot blast.” The “cold blast” is sometimes used. The ores when dug from the mine are generally stamped into powder, then “roasted,”—that is, made hot, and kept so for some time to drive off water, sulphur, or arsenic, which would prevent the “fluxes” acting properly. The fluxes are substances which will mix with, melt, and separate the matters to be got rid of, the chief being charcoal, coke, and limestone. The ore is then mixed with the flux, and the whole raised to a great heat; as the metal is separated it melts, runs to the bottom of the “smelting furnace,” and is drawn off into moulds made of sand; it is thus cast into short thick bars called “pigs,” so we hear of pig-iron, and pig-lead. Iron is smelted from “ironstone,” which is mixed with coke and limestone. The heat required to smelt iron is so very great, that a steam-engine is now generally employed to blow the furnace. (Before the invention of the steam-engine, water-mills were used for the same purpose.) The smelting is conducted in what is called a blast furnace. When the metal has all been “reduced,” or melted, and run down to the bottom of the furnace, a hole is made, out of which it runs into the moulds; this is called “tapping the furnace.”
Smelting is often confounded with melting, as the names are somewhat alike, but the processes are entirely different; in melting, the metal is simply liquefied, in smelting, the metal has to be produced from ores which often have no appearance of containing any, as in the case of ironstone, which looks like brown clay.
The cone of the furnace, A, is lined with fire-bricks, i i, which is encased by a lining, l l; outside are more fire-bricks, and then masonry, m n; C is the throat of the furnace; D the chimney. The lower part, B, is called the boshes. As soon as the ore in the furnace has become ignited the carbon and oxygen unite and form carbonic acid, which escapes, and the metal fuses at last and runs away. The coal and ore are continually added year after year. The glassy scum called “slag “ protects the molten iron from oxidation.
The metal drawn from the blast furnace is “pig iron,” or “cast” iron, and contains carbon. This kind of iron is used for casting operations, and runs into sand-moulds. It contracts very little when cooling. It is hard and brittle.
Bar Iron is the almost pure metal. It is remarkably tenacious, and may be drawn into very fine wire or rolled. But it is not hard enough for tools. It is difficult to fuse, and must be welded by hammering at a red heat. Wire-drawing is performed by taking the metal as a bar, and passing it between rollers (fig. 404), which flattens it, and then between a new set, which form cutting edges on the rolled plate (fig. 405), the projections of one set fitting into the hollows of the other closely as in the illustration. The strips of metal come out at the aperture seen at A in the next illustration. These rods are drawn through a series of diminishing holes in a steel plate, occasionally being heated to keep it soft and ductile. When the wire has got to a certain fineness it is attached to a cylinder and drawn away, at the same time being wound round the cylinder over a small fire. Some metals can be drawn much finer than others. Gold wire can be obtained of a “thickness” (or thinness) of only the 5,000th part of an inch, 550 feet weighing one grain! But platinum has exceeded this marvellous thinness, and wire the 30,000th part of an inch has been produced. Ductility and malleability are not always found in the same metal in proportion. The sizes of wires are gauged by the instrument shown in the margin. The farther the wire will go into the groove the smaller its “size.”
Steel contains a certain amount of carbon, generally about 1 to 2 per cent. Cast steel is prepared from cast iron. Steel from bar-iron has carbon added, and is termed bar-steel. The process is called “cementation,” and is carried on by packing the bars of iron in brick-work boxes, with a mixture of salt and soot, or with charcoal, which is termed “cement.” Steel is really a carbide of iron, and Mr. Bessemer founded his process of making steel by blowing out the excess of carbon from the iron, so that the proper amount—1·5 per cent.—should remain.
A brief summary of the Bessemer process may be interesting. If a bar of steel as soft as iron be made red-hot and plunged into cold water, it will become very hard. If it be then gently heated it will become less hard, and is then fitted for surgical instruments. The various shades of steel are carefully watched,—the change of colour being due to the varying thickness of the oxide; for we know that when light falls upon very thin films of a substance,—soap-bubbles, for instance,—the light reflected from the under and upper surfaces interfere, and cause colour, which varies with the thickness of the film. These colours in steel correspond to different temperatures, and the “temper” of the steel depends upon the temperature it has reached. The following table extracted from Haydn’s “Dictionary of Science” gives the “temper,” the colour, and the uses of the various kinds of steel.
Temperature | Colour. | Uses of Steel. Cent. Fahr.| | ------------+-------------------+------------------------------------- 220° = 430° | Faint yellow | Lancets. 232° = 450° | Pale straw | Best razors and surgical instruments. 243° = 470° | Yellow | Ordinary razors, pen-knives, etc. 254° = 490° | Brown | Small shears, scissors, cold chisels, etc. 265° = 510° | Brown and purple | Axes, pocket-knives, plane-irons, etc. spots 277° = 530° | Purple | Table-knives, etc. 288° = 550° | Light blue | Swords, watch-springs, etc. 293° = 560° | Full blue | Fine saws, daggers, etc. 316° = 600° | Dark blue | Hand and pit saws.
The Bessemer process transfers the metal into a vessel in which there are tubes, through which air is forced, which produces a much greater heat than a bellows does. Thus in the process the carbon of the iron acts as fuel to maintain the fusion, and at the same time by the bubbling of the carbonic acid mixes the molten iron thoroughly.
During the bubbling up of the whole mass of iron, and the extreme elevation of temperature caused by the union of the carbon of the impure iron with the oxygen of the air, the oxide of iron is formed, and as fast as it forms fuses into a sort of glass; this unites with the earthy matters of the “impure” iron, and floats on the upper part as a flux, thus ridding the “cast iron” of all its impurities, with no other fuel than that contained in the metal itself, and in the air used. When the flame issuing from the “converter” contracts and changes its colour, then the time is known to have arrived when the iron is “ de-carbonized.” The amount of carbon necessary is artificially added, ebullition takes place, a flame of carbonic oxide comes out, and the metal is then run into ingots.
The compounds of iron which are soluble in water have a peculiar taste called chalybeate (like ink). Many mineral springs are so flavoured, and taste, as the immortal Samuel Weller put it, “like warm flat-irons.” Iron is frequently used as a medicine to renew the blood globules.
Protoxide of Iron is known only in combination.
Sesqui-Oxide of Iron is “red ironstone.” Powdered it is called English rouge, a pigment not altogether foreign to our use. In a pure state it is a remedy for arsenical poisoning, and is really the “rust” upon iron.
Bisulphide of Iron is iron pyrites, and is crystalline.
Chloride of Iron is dissolved from iron with hydrochloric acid. It is used in medicine.
Cyanide of Iron makes, with cyanide of potassium, the well-known prussiate of potash (ferro-cyanide of potassium), which, when heated, precipitates Prussian blue (cyanogen and iron).
The Sulphate of the Protoxide is known as copperas, or green vitriol, and is applied to the preparation of Prussian blue.
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MANGANESE is found extensively, but not in any large quantities, in one place; iron ore contains it. It is very hard to fuse, and is easily oxidised. The binoxide is used to obtain oxygen, and when treated with potassium and diluted, it becomes the permanganate of potassium, and is used as “Condy’s fluid.” It readily oxides organic matters, and is thus a disinfectant. It crystallizes in long, deep, red needles, which are dissolved in water. It is a standard laboratory test. There are other compounds, but in these pages we need not detail them.
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COBALT and NICKEL occur together. They are hard, brittle, and fusible. The salts of cobalt produce beautiful colours, and the chloride yields an “invisible” or sympathetic ink. The oxide of cobalt forms a blue pigment for staining glass which is called “smalt.” Nickel is chiefly used in the preparation of German silver and electro-plating. The salts of nickel are green. Nickel is difficult to melt, and always is one of the constituents of meteoric iron, which falls from the sky in aërolites. It is magnetic like cobalt, and is extracted from the ore called kupfer-nickel. A small United-States coin is termed a “nickel.”
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COPPER is the next metal we have to notice. It has been known for centuries. It is encountered native in many places. The Cornish copper ore is the copper pyrites. The fumes of the smelting works are very injurious, containing, as they do, arsenic and sulphur. The ground near the mines is usually bare of vegetation in consequence of the “smoke.” Sheet copper is worked into many domestic utensils, and the alloy with zinc, termed Brass, is both useful and ornamental. Red brass is beaten into thin leaves, and is by some supposed to be “gold leaf”; it is used in decorative work. Bronze is also an alloy of copper, as are gun-metal, bell-metal, etc.
Next to silver, copper is the best conductor of electricity we have. It is very hard and tough, yet elastic, and possesses malleability and ductility in a high degree. It forms two oxides, and there are several sulphides; the principal of the latter are found native, and worked as ores. The sulphate of copper is termed blue vitriol, and is used in calico-printing, and from it all the (copper) pigments are derived. It is also used in solution by agriculturists to protect wheat from insects. When copper or its alloys are exposed to air and water, a carbonate of copper forms, which is termed verdigris. All copper salts are poisonous; white of eggs is an excellent remedy in such cases of poisoning.
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LEAD is obtained from galena, a sulphide of lead. It is a soft and easily-worked metal. When freshly cut it has quite a bright appearance, which is quickly tarnished. Silver is often present in lead ore, and is extracted by Pattison’s process, which consists in the adaptation of the knowledge that lead containing silver becomes solid, after melting, at a lower temperature than lead does when pure. Pure lead therefore solidifies sooner.
One great use of lead is for our domestic water-pipes, which remind us in winter of their presence so disagreeably. Shot is made from lead, and bullets are cast from the same metal. Shot-making is very simple, and before the days of breech-loading guns and cartridges, no doubt many readers have cast bullets in the kitchen and run them into the mould over a basin of water or a box of sand. For sporting purposes lead is mixed with arsenic, and when it is melted it is poured through a sort of sieve (as in the cut) at the top of a high tower. (See figs. 413 and 414). The latter illustration gives the section of the shot tower; A is the furnace, B is the tank for melting the lead, and the metal is permitted by the workman at C to run through the sieve in fine streams. As the lead falls it congeals into drops, which are received in water below to cool them. They are, of course, not all round, and must be sorted. This operation is performed by placing them on a board tilted up, and under which are two boxes. The round shot rush over the first holes and drop into the second box, but the uneven ones are caught lagging, and drop into box No. 1. They are accordingly sent to the furnace again.
The next process is to sort the good shot for size. This is done by sieves—one having holes a little larger than the size of shot required. This sieve passes through it all of the right size and smaller, and keeps the bigger ones. Those that have passed this examination are then put into another sieve, which has holes in it a little smaller than the size of shot wanted. This sieve retains the right shot, and lets the smaller sizes pass, and so on. The shot are sized and numbered, glazed by rolling them in a barrel with graphite, and then they are ready for use. Bullets are made by machinery by the thousand, and made up into cartridges with great speed.
The compounds of lead are also poisonous, and produce “colic,” to which painters are subject. Red lead, or minium, is a compound of the protoxide and the binoxide, and may be found native. The former oxide is litharge; white lead, or the carbonate of lead, is a paint, and is easily obtained by passing a stream of carbonic acid into a solution of acetate of lead. It is used as a basis of many paints.
TIN is another well-known metal. It is mentioned by Moses. It possesses a silver-like lustre, and is not liable to be oxidised. The only really important ore is called Tinstone, from which the oxygen is separated, and the metal remains. Cornwall has extensive tin mines. Tin is malleable and ductile, and can be beaten into foil or “silver leaf,” or drawn into wire. It prevents oxidation of iron if the latter be covered with it, and for tinning copper vessels for culinary purposes. The Romans found tin in Cornwall, and the term “Stanneries” was applied to the courts of justice among the tin miners in Edward the First’s time. We have already mentioned the alloys of tin. The oxides of tin, “Stannous” and “Stannic,” are useful to dyers. The latter is the tinstone (SnO_{2}). Sulphide of tin is called “Mosaic gold,” and is much used for decorative purposes.
ZINC is procured from calamine, or carbonate of zinc, and blende, or sulphide of zinc. It has for some years been used for many purposes for which lead was once employed, as it is cheap and light. Zinc is a hard metal of a greyish colour, not easily bent, and rather brittle; but when made nearly red-hot, it can be rolled out into sheets or beaten into form by the hammer. Zinc is about six-and-three-quarter times heavier than water. Like many other metals, it is volatile (when heated to a certain extent it passes off into vapour), and the probable reason that it was not known or used of old is that it was lost in the attempt to smelt its ores. Zinc is now obtained by a sort of distillation; the ores are mixed with the flux in a large earthen crucible or pot.
We have already noticed the alloy of zinc with copper (brass), and the use of zinc to galvanize iron by covering the latter with a coating of zinc in a bath is somewhat analogous to electro-plating. The metal is largely used as the positive element in galvanic batteries, and for the production of hydrogen in the laboratory. Zinc forms one oxide (ZnO), used for zinc-white. The sulphate of zinc is white vitriol, and the chloride of zinc is an “antiseptic.” Certain preparations of the metal are used in medicine as “ointments” or “washes,” and are of use in inflammation of the eyelids.
CHROMIUM. This “metallic element” is almost unknown in the metallic state. But although little known, the beautiful colours of its compounds make it a very interesting study. The very name leads one to expect something different to the other metals—chroma, colour. The metal is procured from what is known as chrome-ironstone, a combination of protoxide of iron and sesqui-oxide of chromium (FeOCr_{2}O_{3}). By ignition with potassium we get chromic acid and chromate of potassium, a yellow salt which is used to make the other compounds of chromium. The metal is by no means easy to fuse.
Sesqui-Oxide of Chromium is a fine green powder employed in painting porcelain.
Chromate of Lead is termed “chrome yellow,” and in its varieties is employed as a paint.
Chromate of Mercury is a beautiful vermilion. There are numerous other combinations which need not be mentioned here.
ANTIMONY was discovered by Basil Valentine. The Latin term is Stibium, hence its symbol, Sb. It is very crystalline, and of a peculiar bluish-white tint. It will take fire at a certain high temperature, and can be used for the manufacture of “Bengal Lights,” with nitre and sulphur in the proportions of antimony “one,” the others two and three respectively.
The compounds of antimony are used in medicine, and are dangerous when taken without advice. They act as emetics if taken in large quantities. Our “tartar emetic” is well known.
Antimony, in alloy with lead and a little tin, form the type metal to which we are indebted for our printing. Type-casting is done by hand, and requires much dexterity. A ladle is dipped into the molten metal, and the mould jerked in to fill it properly, and then the type is removed and the mould shut ready for another type; and a skilful workman can perform these operations five hundred times in an hour,—rather more than eight times a minute,—producing a type each time; this has afterwards to be finished off by others. The metal of which type is made consists of lead and antimony; the antimony hardens it and makes it take a sharper impression. The letters are first cut in steel, and from these “dies” the moulds are made in brass, by stamping, and in these the types are cast.
Stereotype consists of plates of metal taken, by casting, from a forme of type set up for the purpose: an impression was formerly carried on by plaster-of-Paris moulds, but lately what is termed the papier-maché process is adopted. The paper used is now made in England, and the prepared sheet is placed upon the type and beaten upon it. Paste is then filled in where there are blanks, and another and thicker sheet of the prepared paper is placed over all, dried, and pressed. When this is properly done the paper is hardened, and preserves an impression of the type set up. The paper mould is then put into an iron box, and molten metal run in. In a very short time a “stereotype” plate is prepared from the paper, which can be used again if necessary. The metal plate is put on the machine.
There are several compounds of antimony, which, though valuable to chemists, would not be very interesting to the majority of readers. We will therefore at once pass to the Noble Metals.
THE NOBLE METALS.
There are nine metals which rank under the above denomination:—Mercury, Silver, Gold, Platinum, Palladium, Rhodium, Ruthenium, Osmium, Iridium. We will confine ourselves chiefly to the first four on the list.
MERCURY, or QUICKSILVER, is the first of the metals which remain unaltered by exposure to atmospheric air, and thus are supposed to earn their title of nobility. Mercury is familiar to us in our barometers, etc., and is fluid in ordinary temperatures, though one of the heaviest metals we possess. It is principally obtained from native cinnabar, or sulphide of mercury (vermilion), and the process of extraction is very easy. Mercury was known to the ancients, and is sometimes found native. In the mines the evil effects of the contact with mercury are apparent.
This metal forms two oxides,—the black (mercurous) oxide, or suboxide (Hg_{2}O), and the red (mercuric) oxide, or red precipitate. The chlorides are two,—the subchloride, or calomel, and the perchloride, or corrosive sublimate. The sulphides correspond with the oxides; the mercuric sulphide has been mentioned. Its crimson colour is apparent in nature, but the Chinese prepare it in a particularly beautiful form. Many amalgams are made with mercury, which is useful in various ways that will at once occur to the reader.
SILVER is the whitest and most beautiful of metals, and its use for our plate and ornaments is general. It is malleable and ductile, and the best conductor of electricity and heat that we have. It is not unfrequently met with in its native state, but more generally it is found in combination with gold and mercury, or in lead, copper, and antimony ores. The mines of Peru and Mexico, with other Western States of America, are celebrated—Nevada, Colorado, and Utah in particular. The story of the silver mine would be as interesting as any narrative ever printed. The slavery and the death-roll would equal in horror and in its length the terrible records of war or pestilence. We have no opportunity here to follow it, or its kindred metals with which it unites, on the sentimental side; but were the story of silver production written in full, it would be most instructive.
Silver is found with lead (galena), which is then smelted. The lead is volatilized, and the silver remains. It is also extracted by the following process, wherein the silver and golden ore is crushed and washed, and quicksilver, salt, and sulphate of copper added, while heat is applied to the mass. From tank to tank the slime flows, and deposits the metals, which are put into retorts and heated. The mercury flies off; the silver and gold remain in bars.
In some countries, as in Saxony and South America, recourse is had to another process, that of amalgamation, which depends on the easy solubility of silver and other metals in mercury. The ore, after being reduced to a fine powder, is mixed with common salt, and roasted at a low red heat, whereby any sulphide of silver the ore may contain is converted into chloride. The mixture is then placed, with some water and iron filings, in a barrel which revolves round its axis, and the whole agitated for some time, during which process the chloride of silver becomes reduced to the metallic state. A portion of mercury is then introduced, and the agitation continued. The mercury combines with the silver, and the amalgam is then separated by washing. It is afterwards pressed in woollen bags to free it from the greater part of the mercury, and then heated, when the last trace of mercury volatilizes and leaves the silver behind.
Nitrate of Silver is obtained when metallic silver is dissolved in nitric acid. It is known popularly as lunar caustic, and forms the base of “marking inks.” Chloride of silver is altered by light, but the iodide of silver is even more rapidly acted on, and is employed in photography. Fulminating silver is oxide of silver digested in ammonia. It is very dangerous in inexperienced hands. It is also prepared by dissolving silver in nitric acid, and adding alcohol. It cools in crystals. Fulminating mercury is prepared in the same way.
GOLD is the most valuable of all metals,—the “king of metals,” as it was termed by the ancients. It is always found “native,” frequently with silver and copper. Quartz is the rock wherein it occurs. From the disintegration of these rocks the gold sands of rivers are formed, and separated from the sands by “washing.” In Australia and California “nuggets” are picked up of considerable size.
It is a rather soft metal, and, being likewise costly, is never used in an absolutely pure state. Coins and jewellery are all alloyed with copper and silver to give them the requisite hardness and durability. Gold is extremely ductile, and very malleable. One grain of gold may be drawn into a wire five hundred feet in length, and the metal may be beaten into almost transparent leaves 1/200000 of an inch in thickness!
Aqua-regia, a mixture of hydrochloric and nitric acids, is used to dissolve gold, which is solved only by selenic acid, though the free chlorine will dissolve it. Faraday made many experiments as to the relation of gold to light. (See “Phil. Trans.,” 1857, p. 145.) The various uses of gold are so well known that we need not occupy time and space in recording them. Gilding can be accomplished by immersing the articles in a hot solution of chloride of gold and bicarbonate of potash mixed; but the electro process is that now in use, by which the gold precipitates on the article to be plated.
We have already described the process of electro-plating in the case of silvered articles, and we need only mention that electro-gilding is performed very much in the same way. But gilding is also performed in other ways; one of which, the so-called water gilding, is managed as follows. Gilding with the gold-leaf is merely a mechanical operation, but water-gilding is effected by chemistry.
Water-gilding is a process (in which, however, no water is used) for covering the surface of metal with a thin coating of gold; the best metal for water-gilding is either brass, or a mixture of brass and copper. A mixture of gold and mercury, in the proportion of one part of gold to eight of mercury, is made hot over a fire till they have united; it is then put into a bag of chamois-leather, and the superfluous mercury pressed out. What remains is called an “amalgam”; it is soft, and of a greasy nature, so that it can be smeared over any surface with the fingers. The articles to be gilt are made perfectly clean on the surface, and a liquid, made by dissolving mercury in nitric acid (aqua-fortis), is passed over them with a brush made of fine brass wire, called a “scratch-brush.” The mercury immediately adheres to the surface of the metal, making it look like silver; when this is done, a little of the amalgam is rubbed on, and the article evenly covered with it. It is then heated in a charcoal fire till all the mercury evaporates, and the brass is left with a coating of gold, which is very dull, but may be burnished with a steel burnisher and made bright if necessary. In former times articles were inlaid with thin plates of gold, which were placed in hollows made with a graver, and melted in, a little borax being applied between.
When a solution of “chloride of gold” is mixed with ether, the ether takes the gold away from the solution, and may be poured off the top charged with it. This solution, if applied to polished steel by means of a camel-hair pencil, rapidly evaporates, leaving a film of gold adhering to the steel, which, when burnished with any hard substance, has a very elegant appearance. In this way any ornamental design in gold may be produced, but it is not very durable. The gilt ornaments, scrolls, and mottoes on sword-blades, are sometimes done in this way.
PLATINUM is the heaviest of all metals, gold being next. Platinum is practically infusible, and quite indifferent to reagents. It is therefore very useful in certain manufactories, and in the laboratory. It can be dissolved by aqua-regia. The stills for sulphuric acid are made of platinum, and the metal is used for Russian coinage, but must be very difficult to work on account of its infusible property.
In the finely-divided state it forms a gray and very porous mass, which is known as spongy platinum, and possesses the remarkable property of condensing gases within its pores. Hence, when a jet of hydrogen is directed upon a piece of spongy platinum, the heat caused by its condensation suffices to inflame the gas. This singular power has been applied to the construction of a very beautiful apparatus, known as Döbereiner’s lamp, which consists of a glass jar, a, covered by a brass lid, e, which is furnished with a suitable stop-cock, c, and in connection with a small bell jar, f, in which is suspended, by means of a wire, a cylinder of metallic zinc, z. When required for use, the outer jar is two-thirds filled with a mixture of one part sulphuric acid and four parts water, and the stop-cock opened to allow the escape of atmospheric air, the spongy platinum contained in the small brass cylinder, d, being covered by a piece of paper. The stop-cock is then closed, and the bell jar, f, allowed to fill with hydrogen, and after it has been filled and emptied several times, the paper is removed from the platinum and the cock is again opened, when the gas, which escapes first, makes the metal red-hot and finally inflames. This property of platinum is also used in the “Davy” lamp.
The remaining metals do not call for detailed notice.
In conclusion, we may refer to the following statement, which in general terms gives the properties of the metals, their oxides and sulphides for ordinary readers.
GENERAL CLASSIFICATION OF THE METALS.
(Transcriber’s Note: The table has been divided into two parts to fit within width constraints)
The metals admit of being really distinguished by the following table, in which they are presented in several groups, according to their peculiar properties, and each distinguished by a particular name:—
--------------------------------------- | | | METALS. | | | --------------------------------------+ | (A.) Light Metals. | | | |Specific gravity from 0·8 to 1; | | never occur in the uncombined | | state. | --------------------------------------+ | (a.) Alkaline Metals. | | | | 1. Potassium. | | 2. Sodium. | | (Ammonium.) | | | | | | | --------------------------------------+ |(b.) Metals of the Alkaline Earths.| | | | 3. Calcium. | | 4. Barium. | | 5. Strontium. | --------------------------------------+ |(c.) Metals of the Earths proper. | | | | 6. Magnesium. | | 7. Aluminium. | --------------------------------------+ | (B.) Heavy Metals. | | | |Specific gravity from 5 to 21; are | | found chiefly in combination | | with oxygen, and frequently | | with sulphur and arsenic; some | | are native. | --------------------------------------+ | (a.) Common Metals. | | | | Become oxidized in the air. | | | | 8. Iron. 14. Lead. | | 9. Manganese. 15. Tin. | |10. Cobalt. 16. Zinc. | |11. Nickel. 17. Chromium. | |12. Copper. 18. Antimony. | |13. Bismuth. | --------------------------------------+ | (b.) Noble Metals. | | | | Unchangeable in the air. | | | |19. Mercury. 21. Gold. | |20. Silver. 22. Platinum. | ---------------------------------------
-------------------------------------------------------------------------- | Properties of the | |------------------------------------------------------------------------| | Oxides. | Sulphides. | +----------------------------------+-------------------------------------| A|Powerful bases; possessing a |Powerful bases, which oxidize in | | strong affinity for water, and | the air, and form sulphates; | | form with it hydrates. They | when treated with acids evolve | | yield their oxygen to carbon | hydrosulphuric acid. | | only at a white heat. | | +----------------------------------+-------------------------------------| a|Highly caustic; powerful bases, |Caustic; strong bases; very soluble | | separate all other oxides from | in water, and dissolve a | | their combinations with acids; | large quantity of sulphur, which | | are very soluble in water, and | is separated on addition of an | | do not lose their water of | acid as a white powder, termed | | hydration at the highest | milk of sulphur; they were | | temperatures; attract carbonic | formerly termed liver of sulphur.| | acid rapidly from the air. | | +----------------------------------+-------------------------------------| b|Caustic; strong bases; slightly |Caustic; strong bases; dissolve | | soluble in water; lose their | sulphur, and are partly soluble | | water of hydration at a moderate| in water, and partly insoluble. | | heat, and powerfully absorb | | | carbonic acid. | | +----------------------------------+-------------------------------------| c| { Weak bases, |Insoluble in water. | |Feebly caustic. { insoluble in| | |Not caustic. { water. | | | | | +----------------------------------+-------------------------------------| B|Feebler bases than the foregoing, |Neutral compounds; insoluble in | | some are acids; insoluble in | water; antimony and several | | water, and lose their water of | of the rarer metals produce | | hydration at a moderate heat. | compounds with sulphur, which | | | deport themselves as acids. | | | | | | | +----------------------------------+-------------------------------------| a|With few exceptions, are soluble |Those occurring in nature are | | in acids, and, when ignited with| somewhat brass-like in appearance, | | carbon at a red heat, yield | and are termed pyrites | | their oxygen; are, for the most | and blendes. Those which are | | part, fusible and non-volatile. | artificially prepared have | | | peculiarcolours; by heat they | | | are converted into sulphates. | | | | | | | | | | +----------------------------------+-------------------------------------| b|Have more the properties of acids |With the exception of sulphide | | than of bases; are decomposed | of mercury, they leave the pure | | by ignition into oxygen and | metal when ignited. | | metal. | | | | | | | | --------------------------------------------------------------------------
FOOTNOTES:
[21] Requires oxy-hydrogen blow-pipe.
XXX. ORGANIC CHEMISTRY
RADICALS—ACIDS—BASES—NEUTRALS.
In the introduction to these brief chapters upon Chemistry, we said that the science was divided into two sections, the first section consisting of the simple combinations, and the other of compound combinations. The latter being met with chiefly in animal and vegetable matter, as distinguished from dead or inert matter, was termed Organic. This distinction will be seen below.
COMBINATIONS OF SIMPLE GROUPS.
INORGANIC.
I. Elements and their Combinations. (1) Non-Metallic. (2) Metallic.
II. Peculiar Decompositions of the above. (1) By Electricity. (2) By Light.
COMBINATIONS OF COMPOUND GROUPS.
ORGANIC.
I. Compound Radicals and their Combinations.
II. Peculiar Decompositions of the above. (1) Spontaneous. (2) Dry Distillation.
We have already placed before our readers the elements and their simple combinations, and have incidentally mentioned the decomposition by electricity and by light. In the section upon Electricity the positive and negative poles are explained. Oxygen appears always at the positive pole, potassium at the negative. The other simple bodies vary. Chlorine, in combination with oxygen, is evolved at the negative pole, but when with hydrogen at the positive pole. In the series below each element behaves electro-negatively to those following it, and electro-positively to those above it; and the farther they are apart the stronger their opposite affinities are.
ELECTRICAL RELATION OF THE ELEMENTS.
Oxygen. Sulphur. Nitrogen. Chlorine. Bromine. Iodine. Fluorine. Phosphorus. Arsenic. Carbon. Chromium. Boron. Antimony. Silicium. Gold. Platinum. Mercury. Silver. Copper. Bismuth. Lead. Cobalt. Nickel. Iron. Zinc. Hydrogen. Manganese. Aluminium. Magnesium. Calcium. Strontium. Barium. Sodium. Potassium.
The importance of these facts to science is unmistakable, and, indeed, many attempts have been made to explain, from the electrical condition of the elements, the nature of chemical affinity, and of chemical phenomena in general.
Electrotyping is another instance of decomposition by means of electricity, and respecting decomposition by light we know how powerful the action of the sun’s rays are upon plants, and for the evolution of oxygen. The daguerreotype and photographic processes are also instances which we have commented upon. So we can pass directly to the consideration of the compound groups.
In nearly every complex organic compound we have a relatively simple one of great stability, which is termed the radical, which forms, with other bodies, a compound radical.[22] In these complex groups we find certain elements generally,—viz., carbon, hydrogen, nitrogen, sulphur, and phosphorus. Some compounds may consist of two of these, but the majority contain three (hydrogen, oxygen, and carbon). Many have four (carbon, oxygen, hydrogen, and nitrogen), and some more than four, including phosphorus and sulphur. Others, again, may contain chlorine and its relatives, arsenic, etc., in addition. Now we will all admit that in any case in which carbon is present in composition with other simple bodies forming an organic body, and if that body be ignited in the air, it burns and leaves (generally) a black mass. This is a sure test of the presence of carbon, and forms an organic compound. Similarly in decomposition nitrogen and sulphur in combination inform us they are present by the odour they give off. We need not go farther into this question of radicals and compound radicals than to state that a compound radical plays the part of an element in combination. We find in alcohol and ether a certain combination termed Ethyl. This “compound radical” occurs in same proportions in ether, chloride of ethyl, iodide of ethyl, etc., as C_{2}H_{5}; so it really acts as a simple body or element, though it is a compound of carbon and hydrogen. A simple radical is easily understood; it is an element, like potassium, for instance. We may now pass to the organic combinations classified into ACIDS, BASES, and INDIFFERENT, or NEUTRAL, BODIES.
I. ACIDS.
There are several well-known organic acids, which we find in fruits and in plants. They are volatile and non-volatile; acids are sometimes known as “Salts of Hydrogen.” We have a number of acids whose names are familiar to us,—viz., acetic, tartaric, citric, malic, oxalic, tannic, formic, lactic, etc.
Acetic acid (HC_{2}H_{4}O_{2}) is a very important one, and is easily found when vegetable juices, which ferment, are exposed to the air, or when wood and other vegetable matter is subjected to the process of “dry distillation.” Vinegar contains acetic acid, which is distilled from wood, as we shall see presently. Vinegar is made abroad by merely permitting wine to get sour; hence the term Vin-aigre. In England vinegar is made from “wort,” of malt which is fermented for a few days, and then put into casks, the bung-holes of which are left open for several weeks, until the contents have become quite sour. The liquid is then cleared by isinglass. The vinegar of commerce contains about 6 per cent. of pure acetic acid, and some spirit, some colouring matter, and, of course, water. Wood vinegar (pyroligneous acid) is used for pickles. The ordinary vinegar when distilled is called white vinegar, and it may also be obtained from fruits, such as gooseberries or raspberries.
Acetic acid, or “wood vinegar,” is prepared as follows:—There are some large iron cylinders set in brickwork over furnaces, and these cylinders have each a tube leading to a main pipe in which the liquid is received for condensation. The cylinders, which contain about seven or eight hundredweight, are filled with logs of wood, either oak, beech, birch, or ash, the door is closely fastened, and the joints smeared with clay; the fires are now lighted and kept up all day, till the cylinders are red-hot; at night they are allowed to cool. In the morning, the charcoal, into which the wood is now converted, is withdrawn, and a fresh charge supplied; it is then found that about thirty or forty gallons of liquid has condensed in the main tube from each cylinder, the remainder being charcoal and gases which pass off; the liquid is acid, brown, and very offensive, and contains acetic acid, tar, and several other ingredients, among which may be named creosote; it is from this source all the creosote, for the cure of toothache, is obtained. To purify this liquid it is first distilled, and this separates much of the tar; it is then mixed with lime, evaporated to dryness, and heated to expel the remaining tar and other impurities; it is next mixed with sulphate of soda and water, and the whole stirred together; the soda, now in union with the acetic acid, is washed out from the lime and strained quite clear; it is afterwards evaporated till it crystallises, and vitriol (sulphuric acid) then added; finally the acetic acid is distilled over, and the sulphuric acid left in union with the soda, forming sulphate of soda, to be used in a similar process for the next batch of acid. The acetic acid is now quite colourless, transparent, and very sour, possessing a fragrant smell. This is not pure acetic acid, but contains a considerable quantity of water. The acetic acid of commerce, mixed with seven times its bulk of water, forms an acid of about the strength of malt vinegar, perfectly wholesome, and agreeable as a condiment.
Pure acetic acid may be made by mixing dry acetate of potash with oil of vitriol in a retort, and distilling the acetic acid into a very cold receiver; this, when flavoured with various volatile oils, forms the aromatic vinegar sold by druggists. It is a very strong acid, and if applied to the skin will quickly blister it.
Acetate of lead, or sugar of lead, is obtained by dissolving oxide of lead in vinegar. A solution of this salt makes the goulard water so familiar to all. Acetate of lead is highly poisonous.
Acetate of copper is verdigris, and poisonous. Other acetates are used in medicine.
We may pass quickly over some other acids. They are as follows:—
TARTARIC ACID (C_{4}H_{6}O_{6}) is contained in grape juice, and crystallizes in tabular form. The purified powdered salt is cream of Tartar.
CITRIC ACID (C_{6}H_{8}O_{7}) is found native in citrons and lemons, as well as in currants and other fruits. It is an excellent anti-scorbutic.
MALIC ACID (C_{4}H_{6}O_{5}) is found chiefly in apples, as its name denotes (malum, an apple). It is prepared from mountain-ash berries.
OXALIC ACID (C_{2}H_{2}O_{4}). If we heat sugar with nitric acid we shall procure this acid. It is found in sorrel plants.
TANNIC ACID (C_{27}H_{22}O_{17}). It is assumed that all vegetables with an astringent taste contain this acid. Tannin is known for its astringent qualities. The name given to this acid is derived from the fact that it possesses a property of forming an insoluble compound with water, known as leather. Tanning is the term employed. Tannin is found in many vegetable substances, but oak bark is usually employed, being the cheapest. The “pelts,” hides, or skins, have first to be freed from all fat or hair by scraping, and afterwards soaking them in lime and water. Then they are placed in the tan-pit between layers of the bark, water is pumped in, and the hides remain for weeks, occasionally being moved from pit to pit, or relaid, so as to give all an equal proportion of pressure, etc. The longer the leather is tanned—it may be a year—the better it wears.
Skins for gloves and binding are tanned with “sumach,” or alum and salt. Sometimes the leather is split by machinery for fine working. Parchment is prepared from the skins of asses, sheep, goats, and calves, which are cleaned, and rubbed smooth with pumice stone.
Tannic acid, with oxide of iron, produces Ink, for the gall-nut contains a quantity of the acid. All the black inks in use generally are composed of green vitriol (sulphate of iron) in union with some astringent vegetable matter; the best is the gall-nut, although, for cheapness, logwood and oak bark have each been used. An excellent black ink may be made by putting into a gallon stone bottle twelve ounces of bruised galls, six ounces of green vitriol, and six of common gum, and filling up the bottle with rain water; this should be kept three or four weeks before using, shaking the bottle from time to time.
Blue ink has lately been much used; it is made by dissolving newly-formed Prussian blue in a solution of oxalic acid. To make it, dissolve some yellow prussiate of potash in water in one vessel, and some sulphate of iron in another, adding a few drops of nitric acid to the sulphate of iron; now mix the two liquids, and a magnificent blue colour will appear, in the form of a light sediment; this is to be put upon a paper filter, and well washed by pouring over it warm water, and allowing it to run through; a warm solution of oxalic acid should now be mixed with it, and the Prussian blue will dissolve into a bright blue ink.
Red ink is made by boiling chips or raspings of Brazil wood in vinegar, and adding a little alum and gum; it keeps well, and is of a good colour. A red ink of more beautiful appearance, but not so durable, may be made by dissolving a few grains of carmine in two or three teaspoonfuls of spirit of hartshorn.
Marking ink is made by dissolving nitrate of silver in water, and then adding some solution of ammonia, a little gum water, and some Indian ink to colour it. Printers’ ink is made by grinding drying oil with lamp-black.
The powdered gall-nut is an excellent test for iron in water. It will turn violet if any iron be present.
FORMIC ACID (CH_{2}O_{2}) is the caustic means of defence employed by ants, hence the term formic. It can be artificially prepared by distilling a mixture of sugar, binoxide of manganese, and sulphuric acid. On the skin it will raise blisters.
LACTIC ACID (C_{3}H_{6}O_{3}) is present in vegetable and animal substances. Sour whey contains it, and the presence of the acid in the whey accounts for its power of removing from table-linen stains. When what is called “lactic fermentation” occurs, milk is said to be “turned.”
II.
BASES.
The definition of a base is not easy. We have described bases as substances which, combining with acids, form salts, but the definition of a base is as unsatisfactory as that of acid or salt. All vegetable bases contain nitrogen, are usually very bitter, possess no smell or colour, and are insoluble in water. They are usually strong poisons, but very useful in medicine.
The most important are the following bases:—
QUININE is contained in the cinchona (yellow) bark. One hundred parts of the bark have been calculated to yield three of quinine.
MORPHINE is the poisonous base of opium, which is the juice of the poppy, and is prepared chiefly in India and China.
NICOTINE is the active principle of tobacco, and varies in quantity in different tobaccos. Havannah tobacco possesses the least. It is a powerful poison, very oily, volatile, and inflammable.
CONIA is prepared from the hemlock. It is fluid and volatile. It is also a deadly poison, and paralyses the spine directly.
STRYCHNINE is found in poisonous trees, particularly in the nux-vomica seeds of Coromandel. It produces lock-jaw and paralysis. There is no antidote for strychnine; emetics are the only remedy.
The above are chiefly remarkable for their uses in medicine, and in consequence of their highly poisonous character are best left alone by unpractised hands.
A German chemist, named Serturner, was the first to extract the active principle from Opium. The question of opium importation has lately been attracting much attention, and the opinions concerning its use are divided. Probably in moderation, and when used by ordinary people (not demoralized creatures), it does little harm.
Opium is the juice of the “common” poppy, and derives its name from the Greek opos, juice. The plant is cultivated in India, Persia, and Turkey. After the poppy has flowered the natives go round, and with a sharp instrument wound, or puncture, every poppy head. This is done very early in the morning, and under the influence of the sun during the day the juice oozes out. Next morning the drops are scraped off. The juice is then placed in pots, dried, and sent for export. The “construction” of opium is very complicated, for it contains a number of ingredients, the most important being morphia, narcotine, meconic acid, and codeia. It is to the first named constituent that the somnolent effect of opium is due.
III.
INDIFFERENT SUBSTANCES.
There are a great number of so-called “indifferent” substances to which we cannot be indifferent. Such bodies as these have neither acid nor basic properties, and stand no comparison with salts. They are of great importance, forming, as they do, the principal nutriment of animals. Some contain nitrogen, some do not; they may therefore be divided into nitrogenous and non-nitrogenous substances; the former for solid portions of the body, the latter for warmth.
We will take the latter first, and speak of some of them—such as starch, gum, sugar, etc.
STARCH is found in the roots of grain, in the potato, dahlia, artichoke, etc., and by crushing the parts of the plant, and washing them, the starch can be collected as a sediment. In cold water and in spirits of wine starch is insoluble. The various kinds of starch usually take their names from the plants whence they come. Arrowroot is obtained from the West Indian plant Maranta Arundinacea. Cassava and tapioca are from the manioc; sago, from the sago palm; wheat starch, and potato starch are other examples.
If starch be baked in an oven at a temperature of about 300° it becomes, to a great extent, soluble in cold water, forming what is called “British gum”; this is largely used for calico printing and other purposes; if boiled in water under great pressure, so that the temperature can be raised to the same degree, it is also changed into an adhesive sort of gum, “mucilage”; this is the substance made use of by the government officials to spread over the backs of postage and receipt stamps to make them adhere. The starch of grain, during germination, or growth, contains diastase, which converts the starch into gum and sugar; the same effect can be produced by heating starch with diluted sulphuric acid.
GUM found in plants is chiefly procured from the Mimosa trees, from which it flows in drops, and is called Gum Arabic. There are other so-called “gums,” but this is the one generally referred to.
SUGAR exists in fruits, roots, and in the stalks of plants, in the juice of the cane, maple, and beet-root particularly. The canes are crushed, the juice is clarified with lime to prevent fermentation, and the liquid is evaporated. It is then granulated and cleared from the molasses. Sugar, when heated, becomes dark, and is called “caramel.” It is used for colouring brandy, and gives much difficulty to the sugar refiners.
Sugar refining is conducted as follows. The raw (brown) sugar is mixed into a paste with water, and allowed to drain. The sugar thus becomes white. It is then dissolved in water, with animal charcoal and bullocks’ blood. The liquid is boiled, and put into a dark cistern with holes at the bottom, and cotton fibres being fastened in the holes, are hung into another dark cistern, into which the liquid runs pure and white. It is then pumped into a copper vessel,—vacuum pan,—and condensed to the proper consistence. Subsequently it is poured into conical moulds, and pure syrup poured upon the crystal shapes. The caramel is then removed through a hole at the end. The moulds or loaves are then dried, and if not even or elegant they are turned in a lathe. Finally they are packed up as “loaf sugar.” Sugar undergoes no decomposition, and is the cause of non-decomposition in other substances. For this reason it is employed in “preserving” fruit, etc. Sugar is obtained from beet by crushing and rasping the roots, as the cane is treated.
SPIRIT OF WINE, OR ALCOHOL, is not a natural product. It is found by the decomposition of grape-sugar by fermentation. There is a series of alcohols which exhibit a regular gradation, founded, so to speak, upon one, two, or three molecules of water. They are called respectively alcohols, glycols, and glycerins. Thus we have—
Alcohols.
Methylic alcohol. Ethylic ” Prophylic ” Amylic ”
Glycols.
Enthelein glycol. Prophylene ” Butylene ” Amylene ”
Glycerins.
(Ordinary Glycerine is the only one known.)
The cetyl and melissylic alcohols are contained in spermaceti and bees-wax respectively. The usual alcohol is the Vinic, a transparent, colourless liquid, which is the spirituous principle of wine, spirits, and beer, and when sugar is fermented the alcohol and carbonic acid remain.
Spirits of wine has a very powerful affinity for water, and thus the use of stimulants in great quantity is to be deprecated, for alcohol absorbs the water from the mucous membranes of the stomach and the mouth, making them dry and hard. The state of “intoxication,” unfortunately so familiar, is the effect produced by alcohol upon the nerves. We append a list of the beverages which are most in use, and the percentage of alcohol in each according to Professor Hart:—
Port 15 per cent. Madeira 14·5 ” Sherry 14 ” Claret 8 ” Ale 6 ” Porter 5 ”
Spirit of wine is contained in many mixtures, and for the purpose of ascertaining how much alcohol may be in wine, or any other liquid, a hydrometer is used (fig. 432). This instrument consists of a glass tube with a bulb at the end. It is put into water, and the place the water “cuts” is marked by a line on the stem, and called zero 0°. Spirit of wine has less specific gravity than water, so in absolute alcohol the instrument will sink lower than in water, and will descend to a point which is marked 100. In any mixture of alcohol and water, of course the hydrometer will rise or sink between the extreme points accordingly as the mixture may contain less alcohol or more. So a scale can be furnished. The instrument, as described, was invented by MM. Gay-Lussac and Tralles, and called the “percentage” hydrometer. There are many other instruments marked in a more or less arbitrary manner. We append a comparative table of a few hydrometers. (See page 420.)
ETHER, or sulphuric ether, is a mixture of spirits of wine with sulphuric acid, and distilled. It loses water, and the product is ether, which is volatile, and transparent, with a peculiarly penetrating odour. It will not mix with water, and if inhaled will produce a similar effect to chloroform.
COMPARATIVE TABLE OF HYDROMETERS.
---------+------------+------------+-----------+-----------+----------- Specific | Percentage | Percentage | Degree, | Degree, | Degree, Gravity. | Volume | Weight, | according | according | according | (Tralles). | at 60° F. |to Cartier.| to Beck. | to Baumé. ---------+------------+------------+-----------+-----------+----------- 1·000 | 0 | 0 | 10 | 0 | 10 0·991 | 5 | 4·0 | ... | ... | ... 0·985 | 10 | 8·0 | 12 | ... | ... 0·980 | 15 | 12·1 | ... | 3 | 13 0·975 | 20 | 16·2 | ... | ... | ... 0·970 | 25 | 20·4 | 14 | 5 | ... 0·964 | 30 | 24·6 | 15 | 6 | 15 0·958 | 35 | 28·9 | ... | ... | 16 0·951 | 40 | 33·4 | ... | 9 | 17 0·942 | 45 | 37·9 | 18 | ... | ... 0·933 | 50 | 42·5 | ... | 12 | 20 0·923 | 55 | 47·2 | 21 | 14 | ... 0·912 | 60 | 52·2 | ... | 16 | 24 0·901 | 65 | 57·2 | 24 | 19 | ... 0·889 | 70 | 62·5 | 27 | ... | 28 0·876 | 75 | 67·9 | ... | 24 | ... 0·863 | 80 | 73·5 | 30 | 27 | 32 0·848 | 85 | 79·5 | 35 | 30 | 35 0·833 | 90 | 85·7 | ... | 34 | 38 0·815 | 95 | 92·4 | 40 | 38 | 42 0·793 | 100 | 100·0 | 44 | 44 | 48 ---------+------------+------------+-----------+-----------+-----------
Chloroform is transparent, and will sink in water. Diluted alcohol with hypo-chloride of lime, will produce it. When inhaled, chloroform produces a pleasing insensibility to pain, and is useful in surgery.
A certain compound of alcohol with mercury dissolved in nitric acid will cause decomposition, and white crystals will eventuate. These compound crystals are termed fulminating mercury.
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We must now pass rapidly over the few remaining subjects we have to notice, such as fats and soaps, wax, oils, etc.
Fats are of the greatest use to man, particularly in cold climates, for upon them depends the heat of the body. Fatty acid, if liquid, is known as oleic acid; if solid, stearic acid. Soaps are compounds of fatty acids. Many “fats” are consumed as food, others as fuel or for lighting purposes, in the shape of oils. Such oils are not primarily useful for burning. Petroleum and other mineral oils are found in enormous quantities in America.
There are what we term fixed oils, and essential or volatile oils. A list is annexed as given by “Hadyn’s Dictionary of Science”:—
FIXED OILS.
Drying. Linseed oil. Poppy oil. Sunflower oil. Walnut oil. Tobacco-seed oil. Cress-seed oil.
Non-Drying.
Almond oil. Castor oil. Colza oil. Oil of mustard. Rape-seed oil. Olive oil, etc.
ESSENTIAL OILS.
Oil of anise. Oil of bergamot. Oil of carraway. Oil of cassia. Oil of cedar. Oil of cloves. Oil of lavender. Oil of lemon. Oil of mint. Oil of myrrh. Oil of nutmeg. Oil of peppermint. Oil of rose. Oil of turpentine.
Vegetable oils are obtained by crushing seeds; animal oils come from the whale and seal tribe. Paraffin oil comes from coal. Linseed is a very drying oil, and on it depends the drying power of paint. We know olive oil will not dry on exposure to the air. Oiled silk is made with linseed oil. When oil is drying in the air considerable heat is evolved, and if oiled substances be left near others likely to catch fire, spontaneous combustion may ensue. Oil of turpentine is found in the pine and fir trees, and many of the oils above mentioned are used by perfumers, etc., the rose oil, or attar of roses, being an Eastern compound.
Allied to the volatile oils are the RESINS, which are non-conductors of electricity. They are vegetable products. They are soluble in alcohol, in the volatile oils, or in ether, and these solutions are called varnishes; the solvent evaporates and leaves the coating. Turpentine, copal, mastic, shellac, caoutchouc, and gutta-percha are all resinous bodies. Amber is a mineral resin, which was by the ancients supposed to be the “tears of birds” dropped upon the seashore. Moore refers to this in his poetic “Farewell to Araby’s Daughter”—
“Around thee shall glisten the loveliest amber That ever the sorrowing sea-bird has wept.”
Amber is not soluble either in water or alcohol; it is, however, soluble in sulphuric acid. It takes a good polish, and when rubbed is very electrical. It is composed of water, an acid, some oil, and an inflammable gas, which goes off when the amber is distilled.
The well-known camphor is got from a tree called the “Laurus Camphora”; it is a white, waxy substance, and can be obtained by oxidizing certain volatile oils. It is generally produced from the Laurus Camphora in a “still.” The behaviour of a piece of camphor in pure water is curious, but its motions can be at once arrested by touching the water or dropping oil on the surface. This phenomenon is due to the surface tension of the liquid, which diminishes when it is in contact with the vapour of the substance.
NITROGENOUS SUBSTANCES.
There are certain albuminous compounds which we must mention here. These are albumen, fibrine, and caseine. Albumen is the white of egg; fibrine is, when solid, our flesh and muscular fibre, while caseine is the substance of cheese. These are very important compounds, and the albuminous bodies are of the very highest importance as food, for the solid portion of blood, brain, and flesh consist, in a great measure, of them. Albumen, fibrine, and caseine contain carbon, hydrogen, nitrogen, and oxygen, with sulphur and phosphorus.
ALBUMEN. The most familiar and the almost pure form of albumen is in the white of eggs, which is albuminate of sodium. It also exists in the serum of the blood, and therefore it is largely found in the animal kingdom. It can also be extracted from seed or other vegetable substances, but it is essentially the same. Albumen is very useful as an antidote to metallic poisons. It forms about 7 per cent. of human blood. It is soluble up to about 140° Fah.; it then solidifies, and is precipitated in a white mass. Albumen is used in the purification of sugar, etc.
FIBRINE is found in a liquid condition in blood. The vegetable fibrine (gluten) is prepared by kneading wheat flour in a bag till the washings are no longer whitened. Like albumen it is found both in a solid and liquid state.
Caseine is seen in the skin which forms upon milk when heated, and forms about 3 per cent. of milk, where it exists in a soluble state, owing to the presence of alkali; but caseine, like albumen, is only soluble in alkaline solutions. As we have said, it is the principal constituent of cheeses. Caseine is precipitated by the lactic acid of milk, which is produced by keeping the milk too warm. Caseine, or curds, as they are called, are thus precipitated. The milk is said to be “sour,” or turned.
MILK, the food of the young of all mammalia, is composed chiefly of water, a peculiar kind of sugar, butter, and caseine. It is this sugar in milk which causes the lactic acid mentioned above. The actual constituents of milk are as follows:—
Water 873·00 Butter 30·00 Sugar 43·90 Caseine 48·20 Calcium (phosphate) 2·31 Magnesia 0·42 Iron 0·07 Potassium (chloride) 1·44 Sodium 0·24 Soda (with caseine) 0·42 ------- 1000·00
The sugar of milk is non-fermenting, and can be procured from whey by evaporation.
DECOMPOSITION.
We have seen that animals and plants are composed of many different substances, and so it will be at once understood that these substances can be separated from each other, and then the decomposition of the body will be completed. When the sap sinks or dries up in plants they are dead. When our heart ceases to beat and our blood to flow we die, and then, gradually but surely, decay sets in. There is no fuel left to keep the body warm; cold results, and the action of oxygen of the air and light or water decays the body, according to the great and unalterable laws of Nature. “Dust thou art, and unto dust shalt thou return,” is an awful truth. The constituents of our bodies must be resolved again, and the unfailing law of chemical attraction is carried out, whereby the beautiful organism, deprived of the animating principle, seeks to render itself into less complicated groups and their primary elements.
This resolution of the organic bodies is decomposition, or “spontaneous decomposition,” and is called decay, fermentation, or putrefaction, according to circumstances. The Egyptians, by first drying the bodies of the dead (and then embalming them), removed one great source of decay—viz., water, and afterwards, by the addition of spices, managed to arrest putrefaction.
Fermentation is familiar in its results, which may be distilled for spirituous liquors, or merely remain fermented, as beer and wine. Fusel oil is prepared from potatoes, rum from cane sugar, arrack from rice. The power of fermentation exists in nature everywhere, and putrefaction is considered to be owing to the presence of minute germs in the atmosphere, upon which Professors Tyndall and Huxley have discoursed eloquently.
Plants are subjected to a process of decomposition, which has been termed “slow carbonization,” under certain circumstances which exclude the air. The gases are given off, and the carbon remains and increases. Thus we have a kind of moss becoming peat, brown coal, and coal. The immense period during which some beds of coal must have lain in the ground can only be approximately ascertained, but the remains found in the coal-measures have guided geologists in their calculations.
Having already mentioned some products of distillation, we may now close this portion of the subject and pass on to a brief consideration of minerals and crystals. We have left many things unnoticed, which in the limited space at our disposal we could not conveniently include in our sketch of chemistry and chemical phenomena.
FOOTNOTES:
[22] Cyanogen, ethyl, and cacodyl, are compound radicals.