CHEMICAL.
We propose, in this article, to give a short but comprehensive sketch of the methods at present employed by Chemists for the analysis or decomposition of the different substances subjected to their examination, and for ascertaining the constituents of chemical compounds.
This very important branch of the science is quite modern. Bergman was the first who attempted to lay down rules for the analysis of minerals, in his treatise De Mineralorum Docimasia Humida, first published in the year 1780; and, for the analysis of waters, in his treatise De Analyse Aquarum, first published in the year 1778. Many improvements were introduced into every branch of analysis by Bergman, Scheele, and Lavoisier. Klaproth greatly facilitated the analysis of minerals, by introducing the use of caustic potash. He devoted a long and laborious life to this department of chemistry, and reduced every part of it under regular formulas. His Beitrag, in six octavo volumes, published between the years 1795 and 1815, is still the best guide for every person who wishes to become a proficient in this essential department of chemistry. Vauquelin has also devoted a long and active life to chemical analyses. His researches have had a more extensive range than those of Klaproth, for he has not confined himself, like that chemist, to the mineral kingdom, but has devoted much of his time to the analysis of animal and vegetable substances. To him, also, we are indebted for a general formula for the analysis of minerals, which, though not quite adapted to the present state of the science, was of the most essential service when it appeared. It is to it that the writer of this article is in a great measure indebted for his first introduction to the practical knowledge of this subject. To Ross we owe the method of analyzing minerals containing an alkali, which is at present pursued. The introduction of the Atomic Theory formed a new era, in the precision with which chemical analyses were conducted. And Berzelius has given us the most ample collection of precise analytical experiments that have yet appeared. Gay-Lussac's doctrine of volumes has contributed no less essentially to the improvement of the analysis of gaseous bodies. Indeed, it is merely the atomic theory exhibited in another point of view, for there exists a very simple and obvious relation between the specific gravity of a gas and the weight of its atom.
But, in this article, which must, of necessity, be confined to a narrow space, it would be impossible to particularize all the numerous contributors to the improvement of Chemical Analysis. It may be sufficient to mention that the two latest treatises on this subject which we have seen are by Thenard and John. Thenard has devoted the whole of the last volume of his Chemistry, published in 1816, and consisting of 209 octavo pages, to a treatise on the Art of Chemical Analysis. His formulas are rather too brief to be a sufficient guide to those who are quite ignorant of the art. Yet his treatise possesses considerable merit, and was perused by the writer of this article with a good deal of interest, because it made him acquainted with the methods of analysis pursued by the French chemists, and put it in his power to compare them with those which he himself has been in the habit of practising. John's Laboratorium is a much more extensive, and apparently a more complete treatise than that of Thenard. But the writer of this article is not entitled to give an opinion respecting it, as he had it in his possession only for a few hours, and could do nothing more than glance at the general plan.
Chemical substances are very numerous, as they consist of all the bodies to be found on the globe. They may, however, be arranged under certain grand heads; because the mode of analysis followed is analogous for every substance included under the same head; but it differs materially when we proceed from the examination of substances belonging to one head to those of another. The analyses are conducted by means of certain substances which have a chemical action upon the bodies to be subjected to examination. If we were to compose a complete treatise on this subject, it would be necessary to enumerate all these reagents, to describe their properties, and to point out the mode of obtaining them in a state of purity. But this could not be done without occupying a very considerable space; and would, besides, be a repetition of many things which have been already detailed at sufficient length in the body of the Encyclopaedia, and even in some articles of this Supplement. We shall therefore take it for granted, that the readers of this article are already acquainted with the names and the general properties of the different chemical substances which we shall have occasion to mention; and that, if they reduce any of our rules to practice, they will take care to purchase the requisite articles in a state of sufficient purity.
We shall divide this article into eight chapters, in which we shall treat successively of the following classes of bodies: 1. Gases; 2. Salts; 3. Mineral Waters; 4. Metals and their alloys; 5. Stones; 6. Ores; 7. Vegetable Substances; 8. Animal Substances.
CHAP. I.—Of Gases.
The gases at present known, and likely to be the subject of chemical investigation, amount to about 21. Several of these are absorbed by water in such great quantity, that they can be examined only over mercury; but the others not being sensibly absorbed by water, or being absorbed slowly, and in no great quantity, may be examined and prepared over that liquid. We shall divide this chapter into three sections. In the first, we shall take a view of the apparatus which is used by chemists for making experiments on gases; in the second, we shall point out the characters by which the different gases may be distinguished from each other; and, in the third, we shall point out the method of analysing gaseous mixtures, containing three or more gases mixed together in unknown proportions.
Sect. I.—Of the Apparatus.
Before attempting to make experiments on gases, it is necessary to provide both a water-trough and a mercurial trough; the first for receiving those gases that may be examined over water; the second for those that must be examined over mercury. The water-trough may be of any dimensions at pleasure, but it is greatly for the convenience of the experimenter to have at least two; one of a large size, to serve when it is requisite to measure very large volumes of gas, and the other much smaller, in which experiments of research may be conducted. The first may be of wood lined with lead. It may be four feet long, two feet broad, and a foot and a half deep. It should have a shelf about ten inches broad and four feet long, fixed along one of the sides of the trough, about three inches below the surface. This shelf ought to be perforated with three or four holes, varying in diameter from about an inch to a quarter of an inch, and widened out below into a kind of funnel-shape. Over these holes the glass jars destined to receive the gas are to be placed, previously filled with water, and the beak of the retort or glass tube from which the gas issues is made to terminate at the funnel-shaped bottom of the hole. The small trough is best constructed of tinned copper. It may be two feet long, fifteen inches broad, and ten inches deep; and the shelf is best placed at one of its ends. Let it be provided with a single hole near the middle, but towards the edge of the shelf; and it ought to be funnel-shaped below, by soldering a bit of tinned copper round its bottom. This trough may be japanned, which improves its appearance; but the japan soon wears off the inside, if it be much used in experimental investigations. Such troughs are usually constructed of tin-plate; but they wear out so fast, that a chemist will find it cheaper in the end if he get them constructed at first of tinned copper.
One of the most convenient mercurial troughs is the kind made by Mr Newman, of Lisle Street, Leicester Square, London, of which we give a figure, (Plate LXIX. fig. 1.) from the Royal Institution Journal, Vol. I. p. 185. This trough would be much improved by having a round cavity near the end farthest from the glass jar, about eight inches deep and two inches in diameter, or, at least, large enough to admit the common Volta's eudiometer used in this country. The advantage of this hollow would be that the bulk of the gases let into the eudiometer, and the change of bulk produced by the explosion, would be much more conveniently estimated than at present; because the eudiometer tube could be sunk till the mercury within the tube be on a level with the surface of the mercury in the trough. At present, an arithmetical calculation is required to determine the bulk of the gas after the explosion; which, when five or six such experiments are to be made in succession, as is usually the case, makes the labour more tedious and disagreeable than it would otherwise be.
The glass jars destined for holding the gases are usually cylindrical; the larger terminate above in a knob, the smaller are sometimes flat at the top, and sometimes they are hemispherical. They are called test glasses. Fig. 2. represents the shape of the larger glass jars; fig. 3. exhibits a test glass: these last are usually sold in sets, fitted one within the other. The largest of these hold about thirty-five cubic inches, the smallest about two cubic inches. The size of the large glass jars varies from 200 or even 300 cubic inches, down to about 50 cubic inches. All these jars ought to be graduated, in order to indicate the bulk of gas which they contain when at use. The bulks always employed in this country are cubic inches, and tenths, or hundredths of a cubic inch. It is necessary also to be provided with stout glass tubes, as fig. 4, of such a size as to be capable of holding a cubic inch of mercury. These tubes are shut at one end, and rounded off at the other, so that the finger may be applied to shut them up when requisite. They are graduated into 100 equal parts, so that each division is equivalent to the 100th part of a cubic inch.
The method of graduating these tubes being of considerable importance to the practical chemist, deserves to be described before we proceed farther. Provide a tube, as in fig. 5, of a very small bore, and open at both ends, and drawn out at one end into a capillary point. Next take a quantity of clean mercury, and ascertain its specific gravity with all the requisite care. Suppose you find it to be 13.422. Then as at the temperature of 60°, a cubic inch of water weighs 252.72 grains, and \(\frac{4}{10}\)th of a cubic inch 252.72 grains; it follows that at the temperature of 60° the 100th part of a cubic inch of mercury will weigh 33.92 grains. Weigh out 33.92 grains of the mercury, at the temperature of 60°. Put this globule into a watch glass, and plunging the capillary point of the glass tube into the globule, apply your mouth to the other end, and suck the whole of the mercury into the tube. Then apply your finger so as to shut the orifice, and shake the tube till the mercury exactly fills the whole portion of the tube next the capillary extremity. Observe the height of the mercury in the tube, and mark the place with a file, as is done at \(a\) in fig. 5. This tube, supposing it made with sufficient care, will enable us to graduate our measures with considerable expedition. For we have only to plunge its capillary extremity into a quantity of clean mercury in an open vessel, and to suck up the mercury till it reaches the mark \(a\); this quantity (supposing we graduate at the temperature of 60°) is exactly equal in bulk to the 100th part of a cubic inch. Suppose we have a tube to graduate: The first step of the process is to put a narrow ribbon of gummed paper along the whole length of the tube, taking care that the paper be applied as straight as possible. Tie this paper firmly on with a string, which is to be removed when the paper is quite dry. The next step is to place the tube in a perpendicular direction, with its shut end undermost. Take up \(\frac{1}{100}\)th of a cubic inch of mercury, by means of your capillary tube, pour this quantity into the tube to be graduated, apply a small ivory square or rule, so as that its edge shall be a tangent to the convex surface of the mercury in the tube, and draw a fine line, by means of a black lead pencil, across the ribbon of gummed paper. Proceed in this way till you have filled the tube with mercury, and, of course, divided it into the requisite proportional bulks, by means of the lines drawn upon the ribbon of gummed paper. Next take a fine three-cornered file, and wetting the edge of it, draw it carefully along each black line, backwards and forwards, till you have cut through the paper, and made a sufficiently distinct mark upon the tube below it. When the degrees are thus cut, write with a diamond the numbers 10, 20, 30, 40, &c. corresponding to the proper degrees, from the one end of the tube to the other; beginning the divisions, of course, at the shut end of the tube.
The test glasses are graduated exactly in the same manner; but as they are only divided into cubic inches or tenths, a tube holding the requisite quantity of mercury is employed instead of the capillary tube above described. The same method answers for graduating the large glass jars used for holding gases on the water trough; with this difference, that the ten cubic inch measure, employed for the graduation, is filled with water instead of mercury: or they may be placed upon the shelf of the trough, and ten cubic inches of air thrown up. A line is to be made with a black lead pencil, coinciding with the surface of the water in the jar; and this is to be repeated for every ten additional inches, till the whole jar is filled with air. The requisite marks are cut upon the jar by means of a triangular file, and the corresponding numbers, written with a diamond on the glass, precisely as before described.
For many experiments, on certain gases, it is convenient to have test glasses made of thin glass, and bent as in fig. 6. Pieces of phosphorus, sulphur, &c. may be put in the upper part of such tubes, in small platinum trays, and heated by means of a lamp, while surrounded by particular gases, and the changes which take place may be seen through the tube.
The next piece of apparatus which we shall describe is so simple, that it may at first sight appear needless to mention it. But we are of a different opinion, considering it as peculiarly valuable, as it saves a great deal of trouble, and greatly facilitates accurate experimenting with gases over mercury. It is often necessary to transfer determinate quantities of gas from the water trough to the mercurial trough. This must be done without introducing any water along with the gas, except what may exist in the gas in the state of vapour. This is done with the greatest facility by means of the tube represented in fig. 7, which was first used, we believe, for the purpose in question, by Mr Cavendish. This is a glass tube open at both ends; but one of the extremities is bent round, and is drawn out into a fine capillary bore. The tube is graduated into 100th parts of a cubic inch, and the degrees should be made as conspicuous as possible. A good method is to fill the lines, after they have been cut in the glass, with black or red sealing wax. When we want to transfer a given bulk of gas from the water trough to the mercurial trough, we fill this tube with mercury, and shutting the end \(b\) with the finger, we introduce the end \(a\) into a glass jar standing over water, and filled with the gas to be transferred. On removing the finger from the extremity \(b\), the mercury falls down by its weight, into a vessel placed at the bottom of the trough to receive it, and the gas enters by the capillary extremity \(a\) to supply its place. When the gas admitted into the tube amounts to the quantity desired (half a cubic inch for example), we shut the end \(b\) with the finger again. We then withdraw the tube from the water trough, and wipe it dry. The end \(b\) is now to be placed uppermost, and the end \(a\) introduced under the bottom of the glass jar (filled with mercury, and standing over the mercurial
CHEMICAL.
Decomposition, Chemical.
trough) destined to receive the gas. On removing the finger, the weight of the mercury in the tube forces out the gas. When it has all made its way into the destined test-glass, we remove the tube, and we may, in the same way, introduce an additional quantity of any other gas, till we have made up the intended mixture.
Small glass funnels are employed, with points almost capillary, for filling the narrow glass tubes with gas. These funnels, when to be used in the water trough, may be of pewter; but, for the mercurial trough, they must be of glass.
The apparatus for taking the specific gravity of gases is very important; because the specific gravity enters as an essential element in every analysis of gases. It consists of a glass flask, Plate LXIX, fig. 8, thin but stout, and furnished with a stop-cock. The larger and lighter it is so much the better. If its contents be equivalent to fifty cubic inches, it will, in general, be sufficient. The best method is to be supplied with three or four flasks similarly mounted, but differing in capacity from ten cubic inches to fifty. The small ones are used when we are provided with only a small quantity of the gas to be examined, while the larger is preferred when our supply is sufficiently large. We must be furnished with a glass jar similar to fig. 2, only supplied with a stop-cock at the upper extremity. The capacity of this jar must somewhat exceed that of the flask, fig. 8, and the two stop-cocks must be such that they shall screw air-tight into each other.
The process of taking the specific gravity consists of three successive steps. The first step is to weigh the flask, fig. 8, with all the requisite accuracy. We then fix it on the plate of a good air-pump, and exhaust it, and weigh it again. Its weight will be diminished by the quantity of air which has been pumped out of it. In the third place, screw the exhausted flask to the top of the jar containing the gas whose specific gravity is to be ascertained. Great care must be taken that no drops of water have got into either of the stop-cocks. To make matters more certain, the hollow part of the female-screw may be filled with loose cotton wool. When the flask is screwed tight upon the top of the jar, one of the stop-cocks (the lower) is to be opened. You then turn the other stop-cock very slowly round, till you hear the gas beginning to make its way into the flask. Then you stop short. For the smaller the opening is the more slowly does the gas make its way, and the less chance is there of a globule of water being forced in, which would spoil the experiment. When the flask is filled with the gas, lower the jar in the water trough, till the surface of the water, within and without the jar, be on the same level. Then close the stop-cocks,—unscrew the flask, and weigh it again,—noting the change of weight that has taken place. Suppose the weight necessary to be put into the scale to which the exhausted flask is suspended, in order to restore the equilibrium, to be \(a\), and that when the flask is filled with the gas, the weight requisite to be put into the other scale, in order to restore the equilibrium, is \(b\).
It is obvious, that \(a\) represents the weight of the air drawn out of the flask by the air-pump, and \(b\) the weight of a quantity of the gas exactly equal in bulk to the air originally drawn out. Therefore \(a : b : : s : p\); gravity of air; specific gravity of the gas \(= x\). But the specific gravity of air is represented by unity; of course we have this proportion, \(a : b : : 1 : x\), which gives us \(x = \frac{b}{a}\). So that to find the specific gravity of the gas we have only to divide \(b\) by \(a\). Suppose the weight \(a = 15.25\) grains, and \(b = 12\) grains. On such a supposition the specific gravity of the gas would be 0.772.
No correction is necessary, either for temperature or for the height of the barometer, when the experiment is made in the way just described, because all gases undergo the same change in bulk when subjected to the same alteration in temperature or pressure. When the specific gravity of a gas, standing over water, is taken, the gas is always as moist as it can be at the temperature at which the experiment takes place. Thus, if the temperature be 60°, then the vapour in the gas is sufficient to support a column of mercury of the height 0.524 inch. Let us suppose the barometer to stand at 30 inches. On that supposition, \( \frac{1}{4} \)th of the bulk of the gas examined is vapour. The specific gravity of steam, at the temperature of 212°, is 0.625. It is probable that the specific gravity of vapour diminishes with the temperature at which it is produced; but we have not sufficient data to determine at what rate this diminution takes place. If we had, it would be easy to make allowance for the vapour of water present, and to deduce the true specific gravity of the dry gas. But, in the present state of our knowledge, the best way of diminishing the error is, to make the experiment at as low a temperature as possible. Thus, if we take the specific gravity of a gas standing over water, at the temperature of 32°, the elasticity of the vapour which it contains is sufficient only to support a column of mercury 0.2 inch in height; so that the bulk of the vapour present is only about \( \frac{1}{50} \)th of that of the gas under examination. When the gas whose specific gravity is to be taken can only be examined over mercury, the glass jar making a part of the mercurial trough, fig. 1, is to be filled with the gas in question, the flask being such that it can be screwed to this jar as well as to those belonging to the water trough. In such cases, it is easy to dry the gas before taking its specific gravity, by introducing into the jar containing it a little dry muriate of lime, or by opening a communication between the jar holding the gas and a vessel filled with sulphuric acid. This possibility of drying the gases makes it desirable to take, occasionally, the specific gravity even of those that are not absorbed by water over mercury; especially if they happen to be light gases, as hydrogen, because the vapour of water, when it amounts to a sensible quantity, must produce a material difference on the result obtained.
It may be worth while to describe here the mode of taking the specific gravity of the vapours of such liquids as boil at moderate temperatures. The apparatus for this purpose, contrived by Gay-Lussac, is abundantly simple, and answers the purpose sufficiently well. This apparatus is represented in fig. 9. VV is a long narrow glass vessel, capable of holding about 90 cubic inches, and graduated into cubic inches and tenths. It is filled with mercury, and placed inverted in the vessel aa, containing some mercury, and which may be of iron. Small glass globules (as in fig. 10) are blown by the lamp, very thin, nearly globular, but terminating in a capillary tube. One of these glass globules is carefully weighed, and its weight noted down; it is then filled with the liquid, the specific gravity of whose vapour is to be ascertained, and the capillary point is hermetically sealed by the blowpipe. It is then weighed again. The increase in the weight denotes the quantity of liquid, which must be carefully written down. The glass globe is then to be let up into the vessel VV. It will ascend to the top, where it will remain. The cylindrical glass MM, taller than the vessel VV, and open at both ends, is then placed on the outside of VV, and plunged to a certain depth in the mercury of the vessel aa, where it is secured. This cylindrical vessel is now filled with water, and the vessel aa being placed horizontally on a furnace, heat is applied till the water is made to boil. The liquid in the glass globe B expands with the heat, bursts the glass globe, and is converted into vapour. The graduation of the vessel enables us to calculate the bulk of this vapour at the temperature of 212°. Knowing beforehand the weight and specific gravity of the liquid, it is easy to infer the specific gravity of the vapour at 212°. But perhaps it will be of advantage to some readers to give the formula requisite for such calculations, and to exhibit an example of the specific gravity of a vapour deduced from an experiment made by means of this apparatus.
Let P denote the weight of liquid employed, and N the number of divisions of the glass which the same liquid occupies, when converted into a vapour at the temperature of 212°. If we denote the bulk of one of these divisions by r, N will be the bulk of the vapour, on the supposition that the glass does not expand. Of course, if we denote the cubic expansion of glass for one degree by k, 100k will denote the whole expansion of the vessel, on the supposition that we employ the centigrade thermometer, which is more convenient for calculation. The true volume of the vapour will be $N(1+100k)$. This vapour supports the pressure, p, of the atmosphere, minus the height h of the column of mercury still remaining in the glass VV, or it is subjected to the pressure $p-h$. To reduce it to the volume which it would occupy, if it were subjected to a pressure equivalent to a column of 30 inches of mercury, we must multiply $N(1+100k)$ by the inverse ratio of the pressures, 30 and $p-h$, which will change it into
$$\frac{N(1+100k)(p-h)}{30 \text{ inch}}$$
Such is the volume of vapour produced by the weight, P, of the liquid in the circumstances pointed out. If P be expressed in grains, the volume of vapour produced by a grain of the liquid will be
$$\frac{N(1+100k)(p-h)}{P \text{ 30 inch}}$$
When we employ this formula, we must take care to reduce the columns of mercury, p and h, to the same temperature at which the constant pressure of 30 inches is estimated. This is usually the temperature of 32°. The cubic dilatation of common glass for 1 centigrade degree is reckoned 0.0000262716 = k. The expansion of mercury for 1 centigrade degree is $\frac{1}{5412}$.
Such is the formula. Let us now illustrate its application by an example. We shall make choice of an experiment of Gay-Lussac, to determine the specific gravity of steam:
Weight of the little glass globe, 12.2162 grains. Do. filled with water, 21.4826
Weight of water, 9.2664
This portion of water being introduced into the glass vessel, and raised to the boiling temperature, was converted into a quantity of steam, which occupied 220 divisions of the jar. Each of these divisions was equivalent to 0.30472256848 cubic inch. The column of mercury in the glass VV, or h, was in length 2.0473 inches. It was measured by means of the graduated rod T (fig. 9), which passes through the square C.C. This rod is pushed down till it comes in contact with the surface of the mercury in the vessel aa. The moveable piece of wood, H, is then elevated till it correspond with the top of the column of mercury in the vessel VV. It now shows the length of the column upon the rod T.
The height of the barometer was 29.745 inches. The temperature was 15° centigrade.
The first step is to reduce these two columns of mercury to the temperature of 32°. This is done by subtracting from the first
$$\frac{100 \times 2.0473}{5412} = 0.0378,$$
and from the second
$$\frac{15 \times 29.745}{5412} = 0.0824.$$ This subtraction reduces them respectively to 2.0095, and 29.663. We have, therefore,
$$P = 9.2664 \text{ gr.}$$ $$N = 220$$ $$r = 0.3047225 \text{ cubic inch.}$$ $$p = 29.663 \text{ inches.}$$ $$h = 2.0095 \text{ inches.}$$ $$p-h = 27.653.$$
The easiest method of finishing the calculation is by means of a table of logarithms; thus:
| Log. of P | 0.9660110 | |-----------|------------| | Log. 30 | 1.4771212 | | L. P. 30 | 2.4440322 |
| Log. N | 2.3424227 | |--------|------------| | Log. r | 1.4039046 | | Log. p-h | 1.4417423 |
| L. N (p-h) | 3.2680696 | |------------|------------| | L. P. 30 in. | 2.4440322 |
Thus it appears that a grain of water when convert- Decomposition, Chemical.
ed into steam, at the temperature of 212°, and subjected to a pressure of 30 inches of mercury, occupies the bulk of 6.686 cubic inches.
But a grain of water, at its temperature of greatest condensation, is in bulk equal to 0.003953 cubic inch. Hence we see, that when water is converted into steam, its bulk increases
$$\frac{6.686}{0.003953} = 1691.3$$
times. Farther 6.686 cubic inches of steam, at the temperature of 212°, weigh one grain, of course 100 cubic inches weigh 14.95 grains. But 100 cubic inches of common air weigh 30.5 grains. Hence, if we reckon the specific gravity of air, 1, the specific gravity of steam will be 0.4901, at the temperature of 212°, compared to that of air at 60°; or 0.674, if we compare it with air at the temperature of 212°.
When gases, which do not unite chemically with each other, at the ordinary temperature of the atmosphere, are mixed together, no change whatever takes place in their bulk. A cubic inch of hydrogen gas, and a cubic inch of oxygen gas, being mixed together, constitute two cubic inches. Gases thus brought into contact mix together equally. The bulk of each in the above example is doubled, and its elasticity reduced to one half. The hydrogen, instead of one cubic inch, occupies the space of two cubic inches, and so does the oxygen, the density of each, of course, is reduced to one half. So that the specific gravity of such a mixture is exactly the mean of the specific gravity of hydrogen and oxygen gases.
The same case extends to the mixture of vapours, both with gases and with vapours. As long as they continue in the elastic state, they simply mix with gases, without producing any other change than an expansion occasioned by their quantity. If a cubic inch of air, and a cubic inch of any vapour, be mixed together, the mixture will become two cubic inches. The air will expand to twice its former bulk, and so will the vapour, while the elasticity of each will be reduced to one half. Proportional changes will be produced, when vapour is mixed with gases in smaller proportions. This change in bulk is very well illustrated by the piece of apparatus exhibited in fig. 11, which was contrived likewise by M. Gay-Lussac.
AB is a cylindrical glass tube, divided into equal parts, and furnished at its two extremities with the two iron stop-cocks RR'. A little above the lowest stop-cock, another tube of glass TT', communicates with the tube AB. This apparatus being well dried, recently boiled mercury is poured into it by the stop-cock R', till the tube AB is quite filled with that fluid, which will, of course, stand at the height A in the tube TT'. The stop-cock R' being now shut, a globular glass vessel, furnished with the stop-cocks r, containing the gas to be experimented on previously rendered as dry as possible, is to be screwed air-tight upon the stop-cock R'. The stop-cocks r, R' being opened, no gas will enter into the tube AB, provided the gas in the globular vessel be merely subjected to a pressure equivalent to that of the atmosphere. But if we open the stop-cock R, the mercury will run out from its weight. The gas in the globular vessel will expand, and part of it will enter into the tube AB to supply the place of the Decomposed mercury. When a sufficient quantity of the gas has entered the tube AB, the stop-cock R must be shut.
The gas in AB being in a state of expansion, the surface of the column of mercury H will stand higher than the surface of mercury h in the open tube TT', and the difference in height between these two columns will indicate the state of expansion of the gas in AB. To reduce this gas to the same elasticity as that of the external air, we have only to pour mercury into the tube TT', till the surface of the two columns H, h, are precisely upon a level.
To introduce into the tube AB the liquid, the effect of whose vapour on the gas in the tube is to be determined, the stop-cock R'' is to be screwed upon the stop-cock R'. This stop-cock has a very small metallic funnel attached to it, into which the liquid is poured. The cock R'', instead of being perforated, as is usually the case, has merely a depression O cut in it capable of holding a drop of the liquid. This being turned, so that the depression O is in contact with the liquid, a drop of it will fill it, and being now turned round, so that O comes in contact with the stop-cock R', which must be open, the drop of liquid will immediately begin to evaporate, and will in a short time mix itself in the state of vapour with the gas in the tube AB, augmenting its bulk. In this way, as many successive drops of the liquid may be introduced into communication with the gas as are capable of evaporating. We perceive the effect to be at an end, when the mercury ceases to rise any higher in the tube TT'. Let us suppose that a few more drops of liquid have been introduced into the tube than are capable of evaporating. If we open the undermost stop-cock R, and allow the mercury to flow out till the columns H h in the two tubes are precisely upon a level, the pressure to which the gas is subjected in the tube AB, is precisely equal to that of the atmosphere at the time. But its bulk will be greater than it was before the introduction of the vapour.
Let $N =$ volume of the dry gas. $N' =$ volume of do. when mixed with the vapour. $p =$ the height of the barometer. $f =$ the elasticity of the vapour.
It is obvious, that the total elasticity of the mixture of gas and vapour is $f + \frac{pN}{N'}$. But this elasticity is equal to $p$, which we suppose to have remained unaltered during the experiment. Therefore we have $f + \frac{pN}{N'} = p$, and of consequence
$$f = p \cdot \frac{(N-N')}{N'}.$$
In making this experiment, it turns out, that the value of $f$ is always precisely the same as the elastic force of the vapour in a vacuum at the same temperature at which the experiment was made. Hence we may easily calculate beforehand the number $N'$ of divisions which the mixture must occupy under the pressure $p$, supposing the dry gas to have pre-
CHEMICAL.
We have only to ascertain, by means of Dalton's table, the elasticity of the liquid employed ought to have in a vacuum at the temperature of the experiment. Then in the above equation, every thing being known but the number $N'$, it is easy to deduce that quantity, which will be $N' = \frac{p}{p-f} \cdot N$.
Sect. II.—Of the Characters by which the different Gases may be distinguished.
Before it is possible to make experimental researches on gaseous bodies, it is necessary to be familiarly acquainted with the mode of preparing the different gases at present known, with the properties of each of them, and with the means of ascertaining their purity, and distinguishing them from each other. We must in this section confine ourselves to a few observations. But the practical chemist must study in detail and experimentally the history and the properties of every individual gas.
Gases, considered as objects of experiment, must be divided into two classes: 1. Those which may be prepared and examined over water; 2. Those which can only be procured in the gaseous state, by collecting them over mercury. We shall take a view of the characters of the gases belonging to each of these classes separately.
I. Gases permanent over water.
The following is a list of the gases that may be collected and examined over water.
I. Supporters of Combustion. 1. Oxygen. 2. Protoxide of azote. 3. Deutoxide of azote, or nitrous gas. 4. Chlorine.
II. Combustible. 1. Hydrogen. 2. Carbureted hydrogen. 3. Olefiant gas. 4. Protophosphureted hydrogen. 5. Perphosphureted hydrogen. 6. Sulphureted hydrogen. 7. Tellureted hydrogen. 8. Arseniureted hydrogen. 9. Carbonic oxide. 10. Hydrocarbolic oxide.
III. Incombustible. 1. Azote. 2. Carbonic acid.
We shall state the characters by which each of these gases is distinguished, after giving a table of their specific gravity, that of air being reckoned one, and the weight of 100 cubic inches of each at the temperature of 60°, and when the barometer stands at 30 inches.
| GASES | Sp. gravity | Weight of 100 cubic inches | |----------------|-------------|----------------------------| | Air | 1.000 | 30.5 grains | | Oxygen | 1.111 | 33.888 | | Protoxide of azote | 1.528 | 46.598 | | Deutoxide of azote | 1.042 | 31.769 | | Chlorine | 2.500 | 76.25 | | Hydrogen | 0.0694 | 2.117 | | Carbureted hydrogen | 0.555 | 16.99 | | Olefiant | 0.974 | 29.72 | | Protophosphureted hydrogen | 0.972 | 29.634 | | Perphosphureted hydrogen | 0.902 | 27.517 | | Sulphureted hydrogen | 1.180 | 35.89 | | Tellureted hydrogen | | | | Arseniureted hydrogen | | | | Carbonic oxide | 0.972 | 29.652 | | Hydrocarbolic oxide | 0.995 | 30.347 | | Azote | 0.972 | 29.652 | | Carbonic acid | 1.527 | 46.373 |
These three sets of gases may be readily distinguished from each other. If a taper be plunged into a phial filled with any of the supporters, it will burn either more brilliantly than in the open air, as when it is plunged into oxygen gas or protoxide of azote; or its flame will be diminished, and it will give out a great deal of smoke, as when it is plunged into chlorine. When a taper is plunged into deutoxide of azote, it is extinguished. But this gas is at once distinguished by the red colour which it assumes when mixed with common air. When a taper is plunged into a combustible or incombustible gas, it is extinguished. But if the mouth of a phial filled with a combustible gas be held to a lighted candle, the gas will catch fire and burn. Whereas, no such combustion can be perceived when a phial, filled with an incombustible gas, is held to a lighted candle. Thus we can easily determine, whether a gas collected over water be a supporter of combustion, combustible, or incombustible.
1. Oxygen is one of the most important of these gases; because it is employed as a means of ascertaining the nature and constitution of all the combustible gases. Its properties, therefore, ought to be carefully and accurately studied, in the first place, by every person who wishes to become an adept in the analysis of gaseous bodies.
Oxygen gas is colourless,—has no taste nor smell, and is not sensibly absorbed by water. If it be mixed with twice its volume of nitrous gas over water, the mixture becomes red, and is almost wholly absorbed by the water. A solution of protosulphate of iron, recently saturated with deutoxide of azote, if it be put in contact with oxygen gas, in a graduated tube, and the liquid be agitated with it for about five minutes, absorbs the oxygen gas completely. Oxygen gas is absorbed, also, by a solution of hydrogureted sulphuret of lime. It is requisite to have placed this liquid in contact with a certain portion of common air, before it be employed as a test of oxygen; because it has the property of absorbing a certain portion of azote as well as oxygen. On that account we must saturate it with azote beforehand. Unless this precaution be taken, the oxygen under examination might be mixed with a portion of azote without our perceiving it. The simplest and most convenient apparatus for experiments on the purity of oxygen gas by absorption by means of a liquid, is that represented in fig. 12. A is a glass tube capable of holding a cubic inch, and graduated into 100 equal parts. B is a glass cylinder into which the extremity of the tube A fits, being carefully ground to each other with emery. To the glass cylinder B, the caoutchouc bottle C is fixed. The glass cylinder is of such a size as to fit tightly to the mouth of the bottle, and the mouth is tied firmly round the cylinder by a waxed silk thread. The caoutchouc bottle ought to be of such a size as to be capable of containing more than the contents of the whole glass tube A. To examine the purity of oxygen gas by means of this apparatus, we proceed as follows: Fill the tube with the oxygen gas. Fill the caoutchouc bottle with the liquid; for example, a solution of protosulphate of iron, saturated with deutoxide of azote. Plunge the bottle under the surface of the water in the trough, and introduce into its mouth the end of the tube A, which is ground to fit the glass collar B. Remove the apparatus from the trough, and, holding the apex of the tube A downwards, gently squeeze the caoutchouc bottle. The liquid will displace the gas, and gradually fill the tube. The apex of A is now to be turned upwards, and the liquid allowed to run back into the bottle. In this manner the liquid is to be kept running continually backwards and forwards in the tube, till the gas has diminished as much as it is capable of diminishing. The quantity absorbed is shown by the portion of the tube filled with liquid, when the apex of A is held vertical; for the pressure of the atmosphere on the caoutchouc bottle is sufficient to prevent the gas in the tube from remaining in a state of expansion. If the oxygen be pure, it will be absorbed completely. The portion of gas remaining, if there be any, is almost always azote.
Oxygen gas is usually prepared from the peroxide of manganese, or the chlorate of potash. Almost pure oxygen gas may be obtained from manganese, if it be collected after the gas has been coming over for some time, and before the current begins to slacken. The easiest way of determining the purity of oxygen gas is the following: Mix one volume of the oxygen to be examined with two volumes of pure hydrogen gas, and cause an electric spark to pass through the mixture. Note the diminution of bulk. One-third of this diminution is oxygen gas. Let the original volume of the oxygen be \(a\); let the third part of the diminution of bulk produced by the combustion be \(a'\); then the bulk of pure oxygen in the portion of the gas examined is \(a - a'\). Suppose the volume of the oxygen gas examined to be 100, the hydrogen added to be 200, and the diminution of bulk to be 294. The third part of 294 is 98. Here the quantity of pure oxygen in 100 volumes of the gas is 98; of course the oxygen is a mixture of 98 volumes of pure oxygen, and 2 volumes of azote.
The piece of apparatus used for these explosions of oxygen, mixed with any combustible gas, is usually known by the name of Volta's eudiometer. It consists of a very thick glass tube (fig. 13) about eight inches long, and about one-fifth inch of internal diameter, capable of containing about two cubic inches, and graduated into tenths of a cubic inch. Towards the top, two small holes are drilled in it, opposite to each other. Into these, two brass wires are cemented; they penetrate into the inside of the tube, and their blunt ends approach within a small distance of each other. The end of each, on the outside of the tube, terminates in a small ring. A spark from an electrical machine or a Leyden phial, is made to pass from one of these wires to the other within the tube. This spark sets fire to the gaseous mixture, and causes its explosion. The tube, before the explosion, is screwed firmly into an iron claw attached to an iron pillar. This pillar is screwed to the mercurial trough, when the experiments are to be made over mercury; but it is more convenient for experiments over water, to have it screwed to a small moveable box, and to have a little tin plate tube, which is to be filled with water, and the eudiometer put into it before it is fixed to the iron pillar. Fig. 14 represents this iron pillar. A is the short cylinder lined with leather, which opens upon a hinge, and allows the introduction of the glass tube; A is then closed and screwed firmly down. The elasticity of this iron pillar and claw moderates the shock of the explosion, and prevents the eudiometer from being broken. The French chemists have a stop-cock at the bottom of their eudiometers, to shut it during the explosion. The object of this precaution is to prevent the possibility of any of the gas from being driven out of the tube by the expansion which always takes place at the instant of the explosion. But this is sufficiently guarded against by making the experiment upon a small scale. For example, \(\frac{1}{10}\) of a cubic inch of oxygen, and \(\frac{9}{10}\) of a cubic inch of hydrogen gas. The tube should be plunged into a vessel filled with water, and at least six inches deep. When there is a stop-cock at the bottom of the tube, the water of the vessel cannot make its way in to supply the place of the gas condensed by the explosion. The consequence is the formation of a vacuum in the tube. This occasions the evolution of the air, which the water in the tube always contains. This air mixing with the residual gas increases its quantity, and leads the experimenter into an error with respect to the purity of his gas.
2. Protoxide of azote is easily distinguished from oxygen by the following properties: If it be left standing over water, it is gradually absorbed and disappears. For water has the property of absorbing about three-fourths of its bulk of this gas. Hence if a phial be filled with water, and a quantity of this gas be let up into it, if we agitate the phial, a portion of the gas will be absorbed, and the water impregnated with it has acquired a sweet taste. We perceive the same sweet taste if we draw a little of the gas into our mouth through a tube. When a lighted taper is plunged into protoxide of azote it burns with great splendour; but the combustion continues only for an instant or two, much shorter than when the same experiment is made with oxygen gas. The only gaseous body with which protoxide of azote is liable to be contaminated is common air, and azotic gas. It ought to be examined, in the first place, for oxygen gas. This is best done by putting a given volume of it into a jar standing over mercury, and putting up into it a stick of phosphorus with a few drops of water. Allow it to remain in this state for twenty-four hours, in a temperature which should not be lower than 60°. Observe the diminution of the bulk which has taken place. That diminution indicates the quantity of oxygen which the gas contained. If no diminution whatever has taken place, we may conclude that our protoxide of azote contains no oxygen, and, of course, no common air mixed with it.
To determine whether protoxide of azote be contaminated with azotic gas, the best mode of proceeding seems to be the following: Mix a measured portion of the protoxide of azote with its own bulk of pure hydrogen gas, and explode the mixture in a Volta's eudiometer. Note the diminution of bulk. This diminution indicates the quantity, by measure, of protoxide of azote present in the mixture before the explosion. Suppose we mix together 20 measures of protoxide of azote and 20 measures of pure hydrogen gas, and that, after the explosion, we find the residue to be 20 measures, 20 measures of course have disappeared, which indicates 20 measures of pure protoxide of azote. The protoxide in this case was absolutely pure. Suppose a mixture of
- 20 measures protoxide of azote, - 20 measures hydrogen,
and the residue to be 24 measures, 16 measures have disappeared. In this case our protoxide of azote is a mixture of
- 16 measures pure protoxide of azote, - 4 measures azote;
or it contains the fifth part of its bulk of azotic gas.
This mode of examining the purity of protoxide of azote is founded on the following facts, which have been ascertained by experiment. Protoxide of azote is composed of
\[ \text{1 volume oxygen gas} + \text{2 volumes azotic gas} = \text{condensed into 2 volumes.} \]
Hydrogen gas requires half its bulk of oxygen gas in order to be completely changed into water by combustion. Hence, when two volumes of hydrogen are mixed with two volumes of protoxide of azote, the whole hydrogen and oxygen disappears, and there remain two volumes of azotic gas, which is precisely equal to the original bulk of the protoxide of azote.
3. Deutoxide of azote, or nitrous gas, is the most easily recognized of all the gases, and its purity ascertained with great facility. It is colourless and tasteless, and not sensibly absorbed by water. When mixed with common air or oxygen gas, its bulk diminishes, and it assumes a red colour, and the smell of nitric acid. If the mixture be made over water, the red colour disappears very speedily, because the nitric acid formed is absorbed by that liquid. A saturated solution of protosulphate of iron in water absorbs this gas, and acquires a dark brown or almost black colour. To determine the purity of this Decomposition, Chemical gas, fill with it a graduated glass tube, of so narrow a bore that it may be shut up by the point of the finger. Cover its open end with the finger, and introduce it into a small cup, containing a solution of protosulphate of iron. When the liquid has entered perceptibly into the tube, cover its end with the finger, and invert it, that the liquid may run down to the farther extremity of the tube; then introduce it again into the protosulphate, and withdraw the finger. An additional quantity of protosulphate of iron will rush into the tube, to supply the place of the protoxide of azote absorbed. Repeat this process till the bulk of the gas ceases to diminish. The residue indicates the volume of azotic gas with which the nitrous gas is mixed.
4. Chlorine gas has a yellow colour, and a very strong suffocating odour, similar to that of aqua regia. Water absorbs about twice its bulk of it. Hence it may seem, at first sight, an improper gas for being examined over water. But mercury absorbs it very rapidly, and in still greater abundance, forming with it a concrete substance, which lines the inside of the glass jars, and prevents us from seeing what is going on within: so that, upon the whole, it is better to collect this gas over water than over mercury. The best way, when it is prepared for use, is to fill glass phials with it, having ground stoppers, very well fitted, by means of emery.
To determine the purity of chlorine gas, mix it over water with its own bulk of pure hydrogen gas, and explode the mixture in a Volta's eudiometer. Note the diminution of bulk; half of that quantity indicates the bulk of pure chlorine gas in the mixture. The reason of this is, that chlorine combines with its own volume of hydrogen, and is converted into muriatic acid gas, which is absorbed by the water, and disappears.
5. Hydrogen gas, like oxygen, must be familiarly known to the practical chemist, because it is employed to ascertain the purity of the gaseous supporters. We must, therefore, know the methods of obtaining it in a state of sufficient purity, and of determining the foreign gaseous bodies with which it is liable to be contaminated. To prepare hydrogen gas in a state of absolute purity, is a task so difficult, that it seems to have been seldom accomplished. Hence the great difficulty of determining its specific gravity, and the reason why this gravity, as determined by experiment, is somewhat above the truth. But, to procure it pure enough for all the purposes of the examination of gases is sufficiently easy. The most convenient apparatus for the purpose is represented by fig. 15. A is a glass flask, capable of holding about 30 cubic inches; and B is a bent glass tube, fitted by grinding to the mouth of the flask A. Put into the flask a quantity of zinc, broken down into small pieces; then fill the flask to the brim with a mixture of one part of sulphuric acid and three parts of water. Plunge the flask and the bent tube B under the surface of the water, in the water trough; insert the tube B into the flask, taking care, beforehand, that the whole of the air which the tube contained is expelled; now place the flask A upon a stand, so that the end of the bent tube B is under the surface of the water in the water trough, and so situated that a glass jar can be placed over it. The hydrogen gas passes through the bent tube B, and is collected in the glass jar sufficiently pure for all the purposes of experiment.
Hydrogen gas, when pure, has no smell, and, when burnt, gives out so little light, that the colour of the flame cannot be determined. But, as usually prepared, it has a peculiar smell, somewhat similar to that produced by electricity, and it burns with a bluish, greenish, or reddish flame, owing to its being contaminated with phosphorus, sulphur, zinc, or iron. Indeed, when hydrogen is made to issue through a brass tube, and kindled, the colour of the flame is owing to small particles of the brass, which are separated and mix with the flame. Hydrogen is not sensibly absorbed by water. If it be mixed with half its bulk of oxygen gas, and an electrical spark passed through it, an explosion takes place, and the whole gaseous mixture disappears, being converted into water. Therefore, to ascertain the purity of hydrogen gas, mix it with its own volume of oxygen gas, and explode it by electricity. Note the diminution of bulk. Two-thirds of that diminution indicate the bulk of the hydrogen gas contained in the mixture before combustion. Suppose we mix 20 volumes of hydrogen gas with 20 volumes of oxygen gas, and that the diminution of bulk, after the explosion, amounts to 30 volumes. Two-thirds of 30 being 20, we infer from the result, that the hydrogen gas was perfectly pure. Suppose the diminution of bulk to amount to 27 volumes; two-thirds of 27 being 18, we infer that the hydrogen gas is a mixture of
18 volumes pure hydrogen, 2 volumes of some other incombustible gas.
6. Carbureted hydrogen gas is colourless, and burns with a yellowish white flame, like that of a common candle. It is not sensibly absorbed by water, and has neither taste nor smell. It requires for complete combustion twice its volume of oxygen gas, and there remains after combustion a quantity of carbonic acid gas, exactly equal in volume to the original gas. Hence it is obvious, that carbureted hydrogen is a compound of one volume of carbon and two volumes of hydrogen gas, condensed into one volume. One half of the oxygen combines with the carbon, and forms carbonic acid gas; while the other half of the oxygen combines with the two volumes of hydrogen, and is converted into water.
7. Olefiant gas is colourless, and destitute of taste or smell. It burns with a white coloured flame, and gives out much more light than any other gas. Water absorbs about 4th of its weight of this gas; a proportion so small, that it is not perceptible in common experiments. For complete combustion it requires three times its volume of oxygen gas. The residue is carbonic acid gas, constituting twice the volume of the original gas. Hence it is obvious, that olefiant gas is composed of two volumes of carbon and two volumes of hydrogen gas, condensed into one volume. Two volumes of the oxygen combine with the carbon, and constitute two volumes of carbonic acid gas; while the remaining volume of oxygen gas unites with the two volumes of hydrogen, and is converted into water.
When olefiant gas is mixed with its own volume of chlorine, the two gases combine and condense into a liquid, which, if the experiment be made over water, has a white colour, and a considerable opacity. This property is of considerable importance, as it enables us to detect the presence of olefiant gas, when mixed with carbureted hydrogen, carbonic oxide, or any other similar combustible gases. We have only to mix a determinate volume of the gas with chlorine, to allow the mixture to remain for a few minutes over water. If any globules of chloride of olefiant gas appear, we are certain of the presence of olefiant gas in the mixture. By washing out the unabsorbed chlorine, and noting the diminution of bulk, we can determine the proportion of olefiant gas which our mixed gas contains.
8. The two species of phosphureted hydrogen gas may be noticed together. Protophosphureted hydrogen gas may be obtained, by heating crystallized phosphorous acid, which is a hydrate. Perphosphureted hydrogen is procured pure, by putting phosphuret of lime into a small flask, completely filled with water, and receiving the gas which is extracted in glass jars, filled with water, and standing inverted on the shelf of a water trough. Protophosphureted hydrogen is a compound of one volume of phosphorus and two volumes of hydrogen condensed into one volume, while perphosphureted hydrogen is a compound of one volume of phosphorus and one volume of hydrogen, condensed into one volume; or the first may be considered as a compound of one atom phosphorus, and two atoms hydrogen, while the second is a compound of one atom phosphorus and one atom hydrogen. Water absorbs about the eighth part of its volume of each of these gases. They have the smell of garlic. Perphosphureted hydrogen burns spontaneously when it comes in contact with common air or oxygen gas; but perphosphureted hydrogen does not take fire unless it be heated at least to the temperature of 212°, or unless an electric spark is made to pass through it. Perphosphureted hydrogen is absorbed completely by a solution of chloride of lime (oxymuriate of lime). This property, discovered by Mr Dalton, enables us to ascertain the purity of this gas, and to separate it from any other gases with which it may be mixed. When one volume of perphosphureted hydrogen is mixed with one volume, or one and a half volume, or two volumes of oxygen gas over water, combustion takes place in all of these cases, and the two gases completely disappear; the phosphorus being converted into hypophosphorous acid, phosphorous acid, or phosphoric acid respectively, according to the volume of oxygen gas employed.
9. Sulphureted hydrogen gas is colourless. It has a strong disagreeable smell, similar to that of rotten eggs. Water absorbs about two and a half times its volume of this gas when pure; but if the gas be mixed with hydrogen gas, which is very frequently the case, water is scarcely capable of absorbing more than its own bulk of it. This gas burns with a blue flame, and sulphur is deposited during the combustion. If a drop of acetate of lead be let fall into water impregnated with it, a dark brown or blackish precipitate immediately appears. Arsenic is thrown down by it, yellow, and antimony, orange. For complete combustion, it requires one and a half times its volume of oxygen gas; and if the experiment be made over water, the whole gaseous matter disappears, being converted into water and sulphurous acid. Hence it is easy to determine the purity of sulphureted hydrogen gas.
10. Tellurated hydrogen and arseniuretted hydrogen gases are still so imperfectly known, that we cannot attempt to give their characters here. But they are not likely to occur in any analytical investigations of gases.
11. Carbonic oxide gas is colourless, and destitute of taste and smell. It is not sensibly absorbed by water; nor are we acquainted with any substance capable of absorbing it in considerable quantities. It burns with a blue flame, gives out but little light, and is very easily extinguished. For complete combustion, it requires half its volume of oxygen gas. The residue, after combustion, is carbonic acid, and is equal to the bulk of the carbonic oxide before the combustion. If equal volumes of carbonic oxide and chlorine gases, both previously well dried, be mixed together over mercury, and exposed for a short time to the light of the sun, the bulk of the mixture is reduced to one half. If a little water be let up to this new gas, the bulk is not altered; but the gas remaining, which is equal in volume to the original quantity of carbonic oxide employed, is carbonic acid gas, which is absorbed completely, if a sufficient quantity of caustic potash be let up to it.
12. Hydro-carbonic oxide is a gas lately discovered by Dr Thomson, and is by no means unlikely to occur in the complex inflammable gases obtained by distilling vegetable or animal substances. It is easily procured pure by exposing a mixture of prussiate of potash and concentrated sulphuric acid to the heat of a lamp in a small retort. This gas is colourless, has a peculiar smell and an aromatic taste, and is not sensibly absorbed by water. It burns with a deep blue flame, and is much more combustible than carbonic oxide. For complete combustion, three volumes of the gas require to be mixed with two volumes of oxygen gas, and an electric spark passed through the mixture. The residue, which amounts to three volumes, is carbonic acid gas. Hence it is obvious that hydro-carbonic oxide is a compound of three volumes carbonic oxide, and one volume hydrogen condensed into three volumes; one and a half volume of the oxygen unites with the carbonic oxide, and converts it into carbonic acid, without altering its bulk, while the remaining half volume of oxygen unites to the volume of hydrogen, and becomes water.
13. Azotic gas is colourless and destitute of taste and smell, and not sensibly absorbed by water. It extinguishes a burning taper, and is itself incombustible. We are not acquainted with any substance capable of absorbing it, or of combining with it rapidly; so that the properties by which it is distinguished are all negative. If we have a sufficient quantity of this gas to determine its specific gravity, and if we find the negative properties above mentioned to belong to it, we cannot be much deceived in our conclusion respecting the nature of this gas. When very great accuracy is necessary, we may make the following experiment: Mix a small quantity of the gas, supposed to be azote, with twice its volume of oxygen. Let up a little of this mixture into a small glass syphon filled with mercury, and placed so that one leg of the syphon stands in one wine glass containing mercury, and the other in another. Let up to the gas a few drops of water, tinged blue with litmus. Then cause electrical sparks to pass through the gas at the top of the syphon, from one of the columns of mercury to the other. If the gas be azote, the blue liquid will become gradually red, in consequence of the formation of nitrous acid.
14. Carbonic acid gas is colourless. It affects the nose with a peculiar pungent sensation, and has an acidulous and agreeable taste. Water absorbs rather more than its own bulk of it, and hence, if it be allowed to stand upon the water-trough, it soon disappears. It renders lime water, barytes water, and strontian water, milky, and is completely absorbed by a sufficient quantity of these liquids. Hence its nature is easily recognized, and it may be separated from those gases with which it is mixed, and its bulk ascertained. When gases standing over mercury contain a mixture of carbonic acid gas, the best substance to separate it from the other gases with which it is mixed, is a solution of caustic potash in water.
Carbonic acid gas is not only incombustible, but it very much weakens the combustibility of those inflammable gases with which it may be mixed. When carbonic oxide is mixed with a fifth or a sixth of its bulk of carbonic acid gas, it is very difficult to make it burn. When common air is mixed with 1/6th of its bulk of carbonic acid gas, a burning taper, when plunged into it, is immediately extinguished.
II. Gaseous Bodies which must be examined over Mercury.
The following table exhibits a list of these gases.
I. Combustible. 1. Ammonia. 2. Cyanogen.
II. Acid and incombustible. 1. Muriatic acid. 2. Sulphurous acid. 3. Fluoboric acid. 4. Fluosilicic acid. 5. Hydriodic acid.
Before stating the characters of these gases, we shall give a table of their specific gravities, and of the weight of 100 cubic inches of each when the barometer stands at 30 inches, and the thermometer at 60°. The specific gravity of common air is reckoned 1. These two sets of gases may be easily distinguished from each other, by the combustibility of the first set, and the incombustibility of the second.
1. Ammoniacal gas is colourless. Its smell is exceedingly strong and well known, as the odour which distinguishes hartshorn, or the volatile alkali of apothecaries. Its taste is caustic, like that of potash; and, when drawn into the mouth, it very speedily destroys the skin. When placed in contact with water tinged blue with red cabbage, radishes, or violets, it immediately changes the blue colour into green. Water absorbs 780 times its volume of this gas. It is absorbed, likewise, by caoutchouc, and by many other bodies. It is absorbed in considerable quantity by fused muriate of lime, and may be afterwards expelled unaltered from that substance, by the application of a moderate heat.
When ammoniacal gas is put into a dry glass tube, and electrical sparks are passed through it for a considerable time, its bulk is just doubled, and it is converted into a mixture of three volumes hydrogen and one volume azotic gas. Hence it is obvious that ammonia is composed of
\[ \begin{align*} 3 \text{ volumes hydrogen} \\ 1 \text{ volume azote} \end{align*} \]
condensed into 2 volumes.
If we mix two volumes of ammonia with \( \frac{1}{2} \) volume of oxygen, and pass an electrical spark through the mixture, an explosion takes place, water is formed, and there remains a quantity of azotic gas equal to half the original volume of the ammonia. When this gas is mixed with its own volume of sulphurous acid, muriatic acid, or carbonic acid gases, the whole gaseous matter disappears, and a white powder comes in place of it, composed of sulphate of ammonia, muriate of ammonia, or carbonate of ammonia, according to the nature of the acid gas used.
2. Cyanogen gas is colourless. It has a peculiar smell, so pungent that it occasions a flow of tears, and a considerable degree of pain, when applied to the nostrils. This gas burns with a purple flame. Water absorbs \( 4\frac{1}{2} \) times its volume of it, and alcohol 23 times its volume. It reddens the tincture of litmus. If we mix one volume of cyanogen with two volumes of oxygen, and pass an electric spark through the mixture, a violent detonation takes place, and the bulk of the gaseous mixture is not altered; but it is converted into a mixture consisting of two volumes of carbonic acid, and 1 volume of azotic gas. Hence we see that this gas is composed of two volumes carbon and one volume azotic, condensed into one volume. The oxygen combines with the carbon, and is converted into carbonic acid, while the cyanogen is disengaged in the gaseous state.
3. Muriatic acid gas is colourless. It has a peculiar smell, and a strong acid taste. It reddens vegetable blues, and smokes when mixed with atmospheric air, or any gas containing aqueous vapours. Water absorbs 516 times its bulk of this gas, and acquires the appearance and properties of the common muriatic acid of the shops. When mixed with its own volume of ammoniacal gas, the mixture is condensed into a white powder, which possesses the properties of sal ammoniac. If a lighted taper be plunged into a phial full of muriatic acid gas, it goes out, but the flame, before it is extinguished, assumes a green colour. Several of the hydrates of the metallic oxides absorb muriatic acid gas readily. For example, the hydrate of copper, and the hydrate of iron. Black oxide of manganese, when assisted by heat, converts it into chlorine, obviously by absorbing the hydrogen, which constitutes one of its constituents. If this experiment could be made with precision, the muriatic acid would be reduced to one half of its original bulk.
4. Sulphurous acid gas is colourless. It has an exceedingly pungent and disagreeable odour, precisely similar to that of burning sulphur. It reddens vegetable blues, and gradually destroys the colour altogether. If a quantity of peroxide of lead is placed in contact with this gas, it gradually absorbs it completely, being converted into sulphate of lead. By this method, which was first employed by Berzelius, the purity of sulphurous acid gas may be ascertained, and it may be readily separated from any other gases with which it may happen to be mixed. Water absorbs about 43 times its bulk of this gas. When three volumes of sulphured hydrogen gas and two volumes of sulphurous acid gas are mixed together, they condense each other completely into an orange-coloured solid substance, having an acid taste. This action of these two gases on each other was first observed by Mr Kirwan. Berthollet afterwards made some additional remarks on it. But the subject was first accurately examined by the writer of this article, who has found the result as above-stated. When the solid formed is digested in water, it is gradually changed into common sulphur.
5. Fluoboric acid gas is colourless. Its smell is somewhat similar to that of muriatic acid. Its taste is exceedingly acid. Water is capable of absorbing 700 times its bulk of this gas. The specific gravity of this solution is 1.77. It possesses a certain degree of viscosity like sulphuric acid, and requires a high temperature to cause it to boil. This property enables us to distinguish it from muriatic acid. The two gases may likewise be readily recognized by their specific gravities; that of fluoboric acid being much higher than that of muriatic acid gas. It is absorbed in considerable quantities by sulphuric acid. This gas is by no means likely to occur in analytical researches on the gases.
6. Fluosilicic acid gas is colourless. Its taste is intensely acid, and its smell somewhat similar to that of muriatic acid. When it comes in contact with water it is absorbed, and at the same time a quantity of silica is deposited in a gelatinous state. This character is peculiar to fluosilicic acid gas, and is sufficient to distinguish it from every other gas. When water has absorbed 360 times its bulk of this gas, the quantity of silica deposited is so great that the whole liquid has disappeared, being absorbed by the silica; the absorption of course cannot be carried farther.
7. Hydriodic acid gas is colourless. Its smell is similar to that of muriatic acid, and its taste is exceedingly sour. It is distinguished from all the preceding gases by its great specific gravity, and by its not being capable of standing over mercury without undergoing decomposition. The iodine combines with the mercury, and forms iodide of mercury, under the form of a red powder, while a quantity of hydrogen gas remains, equal in bulk to half the hydriodic acid gas. This gas is not likely to occur in chemical analyses of gases. But if it should, it may be easily recognized by the preceding properties.
Sect. III.—Of the Analysis of Gases.
An accurate knowledge of the properties of the different gases, as detailed in the preceding section, will enable the practical chemist to ascertain the constituents of any mixture of gaseous bodies, that may present itself for his examination. We shall, in this section, give a few examples of the mode followed in analysing such mixtures, and we shall chiefly select actual analyses which we have ourselves performed, as more likely to lead the student to a practical knowledge of the subject than imaginary mixtures, never likely to occur, except to the fancy of the writer.
I. Analysis of common Air.
Common air has been ascertained to be a mixture of 21 volumes of oxygen gas, and 79 volumes of azotic gas. To determine its composition, it is requisite merely to know the volume of oxygen gas which it contains. We may infer, without any sensible error, that the residue is azotic gas; for the volume of carbonic acid gas which common air contains is so small, that it may be neglected without any sensible error. The easiest method of analysing common air, is to mix it with half its volume or more of hydrogen gas, and fire the mixture in a Volta's eudiometer, by means of an electric spark. Note the diminution of bulk. One-third of that diminution is the volume of oxygen gas contained in the common air employed. The following experiments were made on a bottle full of air, collected in St James's Park, by carrying the bottle filled with water, and emptying out the water in the Park.
| Volumes of air | Volumes of hydrogen | Residue after Combustion | Diminution of bulk | Oxygen gas in the air | |---------------|---------------------|-------------------------|-------------------|----------------------| | 100 | 50 | 87 | 63 | 21 | | 100 | 50 | 87 | 63 | 21 | | 100 | 60 | 97 | 63 | 21 | | 100 | 60 | 97 | 63 | 21 |
The same method of analysis will answer for any mixture of oxygen and azotic gas, in any proportion whatever.
2. Analysis of Oxygen Gas mixed with an inflammable Gas.
It is possible that a mixture of oxygen gas with hydrogen gas, carbureted hydrogen gas, carbonic oxide, or sulphured hydrogen gas, may be presented for examination. Such a mixture, indeed, is not likely to occur in any gases extricated by chemical processes, however complicated, but it may be made artificially in the laboratory for some purpose or other. How are we to proceed in determining its composition?
1. The first step in the process is to ascertain the presence of oxygen gas, and the proportion of that gas in the mixture. For this purpose, put a small quantity of it (a cubic inch for instance) into a wide graduated tube filled with water, and inverted over the water trough, and let up into it its own volume of deutoxide of azote. If no red colour nor diminution of bulk is produced, we may conclude that the gas contains no oxygen gas. But if the colour becomes red, and the bulk diminishes, oxygen gas is present. To determine the proportion of oxygen present, the easiest method is to use the eudiometer represented in fig. 12. Fill the tube A with the gas, and the bottle C with a solution of protosulphate of iron, saturated with deutoxide of azote. Introduce the tube A into the glass ring B, and agitate the liquid backwards and forwards in the tube till the whole oxygen gas is absorbed. The ascent of the liquid in the tube will indicate the bulk of the oxygen gas.
2. The second step is to determine the nature of the inflammable gas, and its proportion. If we set fire to the portion which we have deprived of oxygen gas, the colour of the flame will give us some idea of the nature of the inflammable gas. If the colour of the flame be blue, the gas is carbonic oxide; if it be white, or yellowish white, the gas is olefiant gas, or carbureted hydrogen; if it be reddish or greenish, and gives out so little light as to be scarcely visible, the gas is hydrogen. But a very little practical skill will enable the experimenter to dispense with this criterion, which at best is rather ambiguous. Take a given volume of the mixed gas, and add to it as much oxygen as will make the volume of the oxygen equal to that of the inflammable gas. Put this mixture into a Volta's eudiometer, and pass an electrical spark through it. A detonation will take place. Note down the diminution of bulk. Fill a graduated tube with lime water, and let up into this tube the residual gas. If it renders the lime milky, the residual gas contains carbonic acid. Hence the inflammable gas under examination must have been carbonic oxide, carbureted hydrogen, or olefiant gas. Note the diminution of bulk produced by the lime water. This diminution, calculated from the original bulk of the residual gas after the explosion, before it was poured out of the Volta's eudiometer, indicates the volume of carbonic acid formed by the combustion. To the residual gas thus deprived of its carbonic acid, let up a measured volume of nitrous gas. If the bulk of this mixture is not diminished, we infer, that the residual gas contains no oxygen gas. If a diminution take place, the residual gas contains oxygen; and the volume of oxygen may be determined by multiplying the diminution of bulk by $\frac{7}{19}$ or $0.3684$. If no oxygen be contained in the residual gas, we must repeat the experiment, mixing a greater proportion of oxygen with the inflammable gas, and we must continue to increase the proportion of oxygen till we find a portion of it remaining in the residual gas after the explosion. This quantity, subtracted from the volume of oxygen present in the mixture before the explosion, gives us the volume of oxygen gas requisite to consume a given volume of the inflammable gas. This fact, together with the quantity of carbonic acid formed by the explosion, puts it in our power to determine the nature of the inflammable gas. If the oxygen necessary for the complete combustion of the gas be $\frac{1}{2}$ the volume of the combustible gas, and if no carbonic acid gas is evolved by the explosion, then the gas is hydrogen. If the gas requires half its volume of oxygen, and if its own volume of carbonic acid gas is formed, the gas is carbonic oxide. If twice the volume of the gas of oxygen by requisite, and if the quantity of carbonic acid formed be equal to the original volume of the inflammable gas, the gas is carbureted hydrogen. If thrice the volume of oxygen be necessary, and if twice the volume of carbonic acid be formed, then the gas was olefiant gas. The following table of experimental results will give the reader a correct idea of the way of performing these experiments.
| No. of Experiments | Volumes of Inflammable Gas | Volumes of Oxygen | Residue after Explosion | Ditto washed with Lime Water | Nitrous Gas added | Diminution of Bulk | |--------------------|---------------------------|------------------|------------------------|----------------------------|-----------------|-------------------| | 1 | 20 | 20 | 10 | 10 | 20 | 27 | | 2 | 20 | 20 | 30 | 10 | 20 | 27 |
In the first of these experiments, it is obvious that the inflammable gas was pure hydrogen. For no carbonic acid was formed by the explosion. The diminution of bulk, when the residual gas was mixed with twice its volume of nitrous gas, was 27 and $27 + 0.3684 = 10$ very nearly. Hence the whole residual gas was oxygen. Thus we see that the gas, for its complete combustion, required half its volume of oxygen, and that no carbonic acid gas was formed by the explosion. But these characters point out unequivocally hydrogen gas.
In the second experiment, we see that a quantity of carbonic acid was formed equal in volume to the volume of inflammable gas, and that the whole residue, after the absorption of the carbonic acid by the lime water, was oxygen. Therefore the inflammable gas required, for complete combustion, half its volume of oxygen gas, and was converted into its own volume of carbonic acid gas. It was therefore carbonic oxide.
From these experiments we may conclude that the gas examined was olefiant gas. The increase of volume observable in the first experiment, when this gas is exploded with a volume of oxygen gas less than its own, is a remarkable property which characterizes this gas. From the average of the other three experiments, we learn that the gas required almost exactly thrice its volume of oxygen gas for complete combustion, while there were formed twice its volume of carbonic acid gas. Now this is the character of olefiant gas.
| No. of Experiments | Volumes of Inflammable Gas | Volumes of Oxygen | Residue after Combustion | Ditto Washed in Lime Water | Nitrous Gas added | Diminution of Bulk | |--------------------|---------------------------|------------------|------------------------|----------------------------|-----------------|-------------------| | 1 | 21.6 | 44.7 | 22 | 0.4 | | | | 2 | 21.6 | 53.5 | 30.1 | 7.1 | 40 | 19 | | 3 | 21.6 | 44.7 | 22.5 | 0.5 | | | | Average | 100 | 205 | 104 | Carbonic acid formed | | |
From these experiments we see that the gas under examination required, for complete combustion, very nearly twice its volume of oxygen gas, and that very nearly its own volume of carbonic acid gas was formed. Now these properties characterize carbureted hydrogen gas.
3. Analysis of a Mixture of Inflammable Gases.
If we have a mixture of different inflammable gases to analyse, the process to be followed is very similar to that exhibited in the last case, the chief difference lies in the inferences drawn from the experiments,
1. If the gaseous mixture is destitute of smell, it can consist only of hydrogen, olefiant gas, carbureted hydrogen, and carbonic oxide, mixed together. If it has the smell of rotten eggs, it contains sulphureted hydrogen; while the smell of garlic indicates the presence of phosphureted hydrogen gas. Sulphureted hydrogen gas may be separated, and its volume determined, by filling the eudiometrical tube, fig. 12, with the gas, filling the caoutchouc bottle... with a solution of sugar of lead in water acidulated with vinegar, and agitating the liquid backwards and forwards, till the bulk of the gas ceases to be diminished. The elevation of the liquid in the tube A, or the diminution of the volume of the gas, indicates the quantity of sulphured hydrogen gas with which it was mixed.
If the presence of phosphureted hydrogen gas be indicated by the smell, the very same method of analysis is to be followed, only the caoutchouc bottle C is to be filled with a solution of chloride of lime in water, instead of a solution of sugar of lead.
2. Suppose the sulphured hydrogen or phosphureted hydrogen removed, the next step is to determine the composition of the residue. Ascertain by various trials the volume of oxygen necessary to consume completely a given volume of the gas, and the quantity of carbonic acid formed. From the knowledge of these facts, it is not difficult to infer the gases, and the proportions of which the gaseous mixture is composed.
For example, if a volume of the gas require 2 volumes of oxygen gas, and produce 1 volume of carbonic acid, we may conclude it to be a mixture of equal volumes of olefiant gas, and carbureted hydrogen gas. If a volume of it require 1 volume of oxygen gas, and if the produce be 1 volume of carbonic acid gas, the gas is a mixture of equal volumes of olefiant gas, carbureted hydrogen, and carbonic oxide. If a volume require its own bulk of oxygen gas, and produce the third of its bulk of carbonic acid gas, we may consider the gas as a mixture of one volume carbureted hydrogen, and two volumes hydrogen.
When olefiant gas is presumed to be present, we may simplify the analysis somewhat by letting up some pure chlorine gas to a given volume of the gas. Chloride of olefiant gas will be formed and separated. If we then remove the excess of chlorine by agitating the gas with some caustic potash, or still better, with lime-water, and note the diminution of bulk, this diminution will indicate the quantity of olefiant gas which was present before the experiment.
When several combustible gases are mixed together, the precision of the analysis is always very much increased, if we can take the specific gravity of the mixed gas. Indeed, if we know the specific gravity, we can in general predict the volume of oxygen gas which will be requisite for complete combustion, and the volume of carbonic acid which will be formed. Suppose the specific gravity of such a mixture to be 0.7645. This being precisely the mean between the specific gravity of olefiant gas and carbureted hydrogen, we conclude, that the gas in question is a mixture of equal volumes of these two gases. We try in consequence whether it will not require for complete combustion 2 times its volume of oxygen gas, and whether it will not produce 1 times its volume of carbonic acid gas. If we find this conjecture verified, no doubt can remain that our opinion respecting the constitution of the gas is correct.
Let the specific gravity of our gas be c. Let there be two gases of known specific gravities; and let these specific gravities be respectively a and b; to determine whether our gas be a mixture of these two gases, and in what proportions. Let x and y denote the respective proportions of the two gases, whose specific gravities are a and b. From a well-known property of fluids we have \( x : y : : c - b : a - c \), hence \( x = \frac{(c - b) y}{a - c} \). But since \( x + y \) must constitute the volume of the gas, which we may call 100, we have \( x + y = 100 \) and \( x = 100 - y \). This gives us the equation \( \frac{(c - b) y}{a - c} = 100 - y \); from which we deduce \( y = \frac{100 (a - c)}{(c - b) + (a - c)} \).
Suppose we have a gas of the specific gravity 0.8128, which is that of a gas that may be extracted from peat by distillation. Let us see whether this gas may not be a mixture of carbonic oxide, and carbureted hydrogen. The specific gravity of carbonic oxide is 0.9722 = a; that of carbureted hydrogen 0.555 = b. In this case \( c = 0.8128 \).
Hence \( y = \frac{100 (a - c)}{(c - b) + (a - c)} = 38.2 \)
\( x = 100 - y = \frac{61.8}{100} \)
We see that a mixture of 38.2 volumes carbureted hydrogen, and 61.8 volumes of carbonic oxide, will produce a gas of the specific gravity possessed by the gas from peat.
38.2 volumes of carbureted hydrogen will require for complete combustion of oxygen gas 76.4 volumes.
61.8 volumes of carbonic oxide 30.9 volumes.
So that 100 volumes of the gas will require for complete combustion 107.3 volumes of oxygen gas.
The carbonic acid gas formed by the combustion of 100 volumes of such a mixture, would be 100 volumes.
Now these proportions agree sufficiently well with several of the experiments related by Dr Thomson in his paper On the Analysis of Gas from Peat (Nicholson's Journal, XVI. 239), to induce us to conclude, that the gas in question is often a mixture of carbureted hydrogen and carbonic oxide, in the proportions here assigned. But the nature of this gas varies with the kind of peat, and with the degree of heat applied to the distillation, as is obvious from the paper just alluded to; and, in many cases, besides carbureted hydrogen and carbonic oxide, hydrogen gas is likewise present.
It would be easy to extend these examples to a much greater length; but the necessary brevity of a supplementary article, obliges us to confine ourselves within as narrow limits as possible. The preceding details, if properly attended to, will enable the pupil to analyse with considerable accuracy, mixtures of gases which are permanent over water. It will be worth the while of those who wish to acquire dexterity in the analysis of gases, to make artificial mixtures of known gases in known proportions, and then to endeavour to detect their composi- tion by analysis. A few weeks' practice in this way will give him a degree of practical skill, which he could not easily have anticipated, provided all the experiments be conducted with precision. Slovenly experiments, how long so ever continued, never lead to accurate knowledge or to an increase of skill.
4. Analysis of gaseous mixtures, which require to be examined over mercury.
We may have over mercury mixtures of the acid gases with each other, or with carbonic acid, sulphurated hydrogen, or chlorine; or with oxygen, azote, hydrogen, or any of the other inflammable gases. A few observations on these cases, which occur but rarely in practice, will suffice.
1. The five acid gases contained in the table, are all absorbed in great quantity by water. If two of them should happen to be mixed together, we are not in possession of methods of separating them from each other, sulphurous acid excepted, which is readily absorbed by peroxide of lead; while the other acid gases have no action on that substance. Our only resource in such a case would be, to cause water to absorb the gaseous mixture,—to saturate both acids with ammonia, and afterwards to throw down each acid by its proper precipitant, and to determine the amount of each by drying the precipitates and weighing them. It seems scarcely necessary to enter into particulars, as such cases are very unlikely to occur, at least to beginners.
2. But nothing is more likely than to have standing over mercury a mixture of carbonic acid gas with sulphurous acid, or muriatic acid, in unknown proportions. Take a determinate volume of such a mixture, and put up into it pieces of common borax. This sub-salt has the property of absorbing the strong acid gases pretty rapidly; but it does not act sensibly on carbonic acid. It will, therefore, absorb the whole of the muriatic acid, or sulphurous acid, and leave the carbonic acid, and the diminution of bulk will enable us to determine the volume of gas which has been withdrawn. When the absorption is at an end, pour the residual gas into another jar, and let up into it a quantity of lime or barytes water. The water will become milky, and the carbonic acid will be absorbed.
The same method will answer if muriatic acid and sulphureted hydrogen gas should happen to be mixed, only instead of lime water we must employ a solution of sugar of lead to absorb the sulphureted hydrogen gas.
3. If we have a mixture of carbonic acid and chlorine gas, we have only to put a given volume of it into a dry jar over mercury, and leave it for 24 hours. The mercury will absorb the whole of the chlorine and leave the carbonic acid.
4. If we have a mixture of an acid gas with oxygen, azote, hydrogen, or any other gas not sensibly absorbed by water, the best mode of proceeding is, to separate the acid gas by means of a solution of caustic potash. The residual gas may then be transferred to the water trough, and examined by the rules already laid down.
The examination of gasses has been greatly facilitated by Gay-Lussac's doctrine of volumes, first explained in the 2d volume of the Memoires d'Arcueil, and to be found in all the recent Treatises on Chemistry. Every person who wishes to become an expert experimenter on gases, ought to make himself familiarly acquainted with this important branch of chemistry.
CHAP. II.—Of Salts.
The salts constitute perhaps the most important set of substances that engage the attention of the practical chemist. It is by their means chiefly that the different acids, alkalis, earths, and metallic oxides, are distinguished from each other. Hence a knowledge of their character, and of the mode of analysing them, is absolutely necessary. A correct analysis of a salt is attended with more difficulty than will at first sight appear. The processes must be as few as possible, and the filtrations and transvasations no more than are absolutely necessary. The difficult part of the process is to wash, dry, and weigh, the precipitates without any loss. We shall divide this chapter into two sections. In the first section we shall give the characters, by means of which, when a salt is presented for examination, its name and composition may be determined; in the second section, we shall give the formulae for analysing with precision the different genera and species of salts.
Sect. I.—Of the Characters of the Salts.
The salts naturally divide themselves into two great orders, according to the nature of the acid which they contain. For one division of the acids is combustible, while another is incombustible. If a salt is presented to us for examination, the first step towards the discrimination of it is to expose it to heat. If it blackens, or if it gives out an inflammable gas, we may consider the salt as belonging to that order which is composed of combustible acids. But if the salt undergoes no change, or if it is merely dissipated without giving out inflammable gas, the salt belongs to the order which is composed of incombustible acids.* We shall therefore divide the salts into orders, and proceed to give the characters of each in succession.
Order I.—Incombustible Salts.
These salts may be divided into the following genera, deriving their names from the acid which they contain.
| I. Acid gaseous. | I. Acid not gaseous. | |-----------------|---------------------| | 1. Carbonates. | 1. Sulphates. | | 2. Sulphites. | 2. Nitrates. |
* The ammoniacal salts may occasion some ambiguity. But they are easily known by mixing them with caustic potash. Fumes of ammonia are exhaled, at once distinguishable by the smell.
CHEMICAL.
1. If an incombustible salt be presented to us for examination, reduce a little of it to powder, and throw it into sulphuric acid. If an effervescence take place immediately, or upon the application of a slight heat, the salt belongs to the genera in the first column of the preceding table.*
If the gas extricated be colourless, and have no other smell but a slight pungency, the salt is a carbonate.
If the gas has a strong smell of burning sulphur, the salt is a sulphite. If the same smell be perceptible, but a quantity of sulphur is deposited, and makes its appearance in the sulphuric acid, the salt is a hypo-sulphite.
If the gas disengaged is red, with the smell of nitrous acid, the salt is a nitrite.
If the salt assumes an orange colour, and if fumes of chlorine be disengaged, it is a chlorate.
If the gas has the smell of rotten eggs, and if it blackens paper, moistened with a solution of sugar of lead, the salt is a hydrosulphuret.
If the gas has a strong smell, and if it forms a white vapour when it mixes with the air, the salt is either a muriate, fluate, or fluoroborate. If it is absorbed by a small quantity of water, and if the water forms a curdy precipitate, when nitrate of silver is dropped into it, which precipitate is redissolved by an excess of ammonia, the salt is a muriate. If the salt effervesces in a glass vessel, but not in a metallic vessel, and if the glass vessel is evidently corroded, the salt is a fluate. If the gas blackens paper when placed in contact with it, the salt is a fluoroborate.
If the gas is a mixture of sulphurous acid, and vapour of iodine, easily distinguished by its violet colour, the salt is a hydriodate.
2. If no effervescence take place, when the salt is mixed with sulphuric acid, even though the mixture be a little heated, we may conclude that the salt belongs to some one of the genera in the second column of the preceding table.
If nitrate of barytes, dropped into the solution of the salt in water, occasions a white precipitate, what is not redissolved, by adding an excess of nitric acid, we may conclude that the salt is a sulphate. But if any doubt remains respecting the accuracy of this conclusion, mix together in a flask ten parts of water, one part of the salt, and two parts of nitrate of barytes, and boil the mixture for a few minutes. A white precipitate will fall. Collect this precipitate on a filter, wash it well, dry it, mix it with its own weight of charcoal powder, and expose it for two hours to a strong red heat; if it be sulphate of barytes, it will be converted into a sulphuret, distinguishable by its smell of rotten eggs, and by its being partly soluble in water, and by the solution letting fall a copious white powder, which is sulphate of barytes, when nitric acid is poured into it.
A nitrate is distinguished by the white fumes, with the smell of nitric acid, which it exhales when heated along with sulphuric acid, and by the brilliant combustion which is produced when a little of it is thrown upon a red hot coal.
A phosphate is recognized in this manner, 1. If it be soluble in water, sugar of lead dropped into the solution occasions a white precipitate. Wash this precipitate and dry it, and expose a little of it on charcoal to the action of the blow-pipe. It immediately melts, and, on cooling, crystallizes in the form of a garnet dodecahedron. 2. If it be insoluble in water, expose a little of it to the flame of the blow-pipe. If it melts, it will crystallize on cooling into a garnet dodecahedron. If it does not melt before the blow-pipe, it will readily dissolve in muriatic acid, and be again precipitated unaltered by caustic ammonia.
The phosphites, when strongly heated, emit phosphureted hydrogen gas, which burns with a flame and a smell that are easily recognizable.
The hypophosphites are all soluble in water, and, when strongly heated, it is probable that they likewise will exhale phosphureted hydrogen; though the experiment has not been tried, and hitherto these salts are very imperfectly investigated. They are not likely to come in the way of an experimenter.
The iodates are all very little soluble in water. When treated with sulphurous acid, or sulphuretted hydrogen, they are decomposed and iodine deposited, easily distinguishable by its colour, and by the property of assuming the state of a violet-coloured vapour, when exposed to heat.
If a borate be dissolved in water, and sulphuric acid dropped in, till the solution acquires a perceptibly acid taste, and then be set aside for some hours, a quantity of boracic acid is deposited in scales, easily recognizable by its properties. It is always in our power to obtain the borate in a state capable of dissolving in water. If it be insoluble, we must put it in a flask, with thrice its weight of carbonate of soda, and ten times its weight of water, and boil it for some time; then saturate the whole of the soda with acetic acid, evaporate to dryness, and digest the dry mass in alcohol. The acetate of soda will be dissolved, and nothing will remain but borate of soda, which is soluble in water, and may be decomposed by means of sulphuric acid.
If we mix an arseniate or an arsenite, with half its weight of black flux or charcoal powder, put the mixture into a glass tube, and expose it to a red heat in a crucible filled with sand, a quantity of metallic
* Prussiate of potash effervesces with sulphuric acid; but it gives out an inflammable gas. It is easily known by its yellow colour, its cubic or octahedral form, and its toughness. arsenic will be sublimed, and coat the upper part of the tube. This arsenic is easily known by its colour, its lustre, its easy combustibility, with the smell of garlic, and the readiness with which it may be sublimed by heat.
If the salt be an arsenite, it precipitates sulphate of copper green, and is precipitated instantaneously yellow, by hydrosulphet of potash. If it be an arseniate, it precipitates sulphate of copper bluish white, and no immediate change is produced on it by hydrosulphuret of potash.
If the salt be a chromate, its colour is yellow. Its solution is precipitated of a very rich yellow by acetate of lead, violet by the nitrate of silver, and red by the protonitrate of mercury.
If the salt be a molybdate, on dissolving it in water, and pouring sulphuric acid into the solution, molybdic acid precipitates in the state of a fine powder. If after this we plunge a piece of tin into the solution, the liquid gradually assumes a blue colour.
When a tungstate is dissolved in water, and the solution mixed with sulphuric, nitric, or muriatic acid, a white floccy precipitate falls. If this precipitate be boiled with a quantity of the acid employed as the precipitant, it assumes a yellow colour, and becomes pure tungstic acid.
Having thus ascertained the genus of our salt by means of the preceding characters, the next point is to determine the species of the salt. This is done by determining the nature of the base, which is united in it with the acid.
If the salt be soluble in water, and if the solution is not rendered muddy by the addition of potash, soda, ammonia, or their carbonates, or by the hydrosulphurets, we are entitled to conclude that the base is a fixed or volatile alkali. Mix a little of it with quicklime and a little water; if a strong ammoniacal odour be exhaled, the base is ammonia. If no odour be exhaled, the base is potash or soda. To determine which of the two, drop into the concentrated solution, a solution of tartaric acid till the liquid tastes sensibly sour. If the base be potash, a white and pretty copious crystalline precipitate will gradually fall, consisting of bitartrate of potash. Into another portion of the solution, drop some muriate of platinum. If potash be the base, a yellow precipitate of muriate of platinum and potash will fall. If neither tartaric acid nor muriate of platinum occasion any precipitate, we may consider the base of the salt under examination to be soda.
If the salt is soluble in water, but precipitated by an alkali or its carbonate, we must dissolve a portion of it in water, and precipitate it by means of carbonate of potash. The precipitate will consist either of the base of the salt in a state of purity, or of that base combined with carbonic acid.
If the salt is insoluble in water, we must reduce a portion of it to powder, mix it with three or four times its weight of carbonate of potash, and ten or twelve times its weight of water, and boil it for a considerable time in a glass flask. By this method the base is gradually separated, either pure or in combination with carbonic acid.
We must now subject the base thus separated to a chemical examination, in order to determine its nature. To assist the experimenter in drawing his conclusions, we shall state the characters by which these different bases are distinguished.
1. Lime, or its carbonate, is a white powder which dissolves readily in nitric and muriatic acids. This solution does not easily crystallize; but when evaporated, leaves a white matter, easily fusible by heat, and very soluble in water. This solution, if not too much diluted, is precipitated white by sulphate of soda. Though dissolved in 100 times its weight of water, it is copiously precipitated by oxalate of ammonia. The salt formed by uniting lime to nitric or muriatic acid is very soluble in alcohol.
2. Barytes dissolves readily in nitric or muriatic acid. The solutions when concentrated by evaporation yield crystals. The crystals of the nitrate are octahedrons or fragments of octahedrons; those of the muriate are usually four-sided tables. Neither of these salts is soluble in alcohol. When sulphuric acid or a sulphate is poured into a solution of either of these salts in water, a white precipitate falls, which is not redissolved by the addition of nitric acid.
3. Strontian is equally easily soluble in nitric or muriatic acid, and the solutions crystallize with still greater facility than those of barytes in the same acids. The nitrate crystallizes in six-sided tables, with bevelled edges; the muriate in long needles, which, when viewed through a glass, have the form of six-sided prisms. The muriate is soluble in alcohol, and the solution burns with a beautiful red flame. When the solution of nitrate or muriate of strontian is diluted with a great deal of water, sulphate of soda does not occasion an immediate precipitate in it, though it precipitates the same salts of barytes, when equally diluted with water.
4. Magnesia dissolves with great facility in sulphuric, nitric, and muriatic acids. The solution in sulphuric acid crystallizes easily in four-sided prisms with square bases, having an intensely bitter taste; but the nitrate and muriate of magnesia are not easily obtained in a crystallized form. If the crystals of the sulphate of magnesia be dissolved in water, and some sulphate of ammonia be poured into the solution, no change takes place. But if a drop of phosphoric acid be let fall into the mixture, a white insoluble powder immediately falls. Carbonate of magnesia may be completely dissolved by mixing it with water, and causing a current of carbonic acid gas to pass through the liquid.
5. Alumina, when newly precipitated, is easily dissolved by caustic potash held in solution by water. If sal ammoniac be poured into this solution, the alumina is again precipitated in white flocks. Alumina may be dissolved in diluted sulphuric acid by means of heat. If this solution, which is colourless, be mixed with a solution of sulphate of potash, and set aside for some time, transparent octahedral crystals of alum are deposited, readily distinguished by their shape and their taste.
6. Yttria is a white powder, which is not dissolved by digesting it in the caustic alkalies, but it dissolves in liquid carbonate of ammonia. It is easily dissolved in sulphuric, nitric, muriatic, and acetic acids. With these acids it forms crystals, which have a A solution of any of these salts is precipitated white by phosphate of soda, carbonate of soda, oxalate of ammonia, and tartrate of potash. It is precipitated likewise by prussiate of potash.
7. Glucina, like alumina, is soluble in the liquid caustic alkalies, and, like yttria, it dissolves in carbonate of ammonia; but its solubility in that liquid is five times greater than that of yttria. Neither the sulphate nor acetate of glucina crystallize. Glucina is thrown down yellow by the infusion of nut-galls, and white by the prussiate of potash.
8. Zirconia is insoluble in the liquid fixed alkalies; but soluble in the alkaline carbonates. It dissolves in nitric, acetic, and muriatic acids; but not in the sulphuric. The nitrate and acetate do not crystallize. The muriate may be obtained in crystals, and is very soluble in water. The solution of muriate of zirconia is precipitated by prussiate of potash, gallic acid, and infusion of nut-galls.
9. Iron. There are two oxides of iron; the black and the red. The first forms crystallizable salts, with sulphuric, muriatic, and acetic acids. These salts have a green colour, and a sweetish, astringent taste. The red oxide dissolves in sulphuric and muriatic acids; the solutions are reddish brown, have an exceedingly strong astringent taste, and do not crystallize. The oxides of iron dissolve readily in muriatic acid. The solution is either green or yellow, according to the acid. Prussiate of potash strikes a dark blue, and infusion of nut-galls a black or a purple, when dropt into this solution. Ammonia throws down the iron in brown flocks. Phosphate of soda occasions a white precipitate. Succinate of ammonia occasions a flesh-coloured precipitate, provided the iron be in a state of peroxide.
10. Nickel. Though this metal forms two oxides, yet as only one of them is found united with acids, constituting salts, we may neglect the other here, which is the peroxide, and has a black colour. All the salts containing the protoxide of nickel are of a fine green colour. The oxide itself is grey; but its carbonate (at least while in the state of hydrate) has a light dirty green colour. This oxide (and the observation applies to most metallic oxides) is not soluble in acids without great difficulty, if it has been exposed to a red heat. The hydrated oxide, or still better, the carbonate, on the contrary, dissolves with facility in sulphuric, nitric, muriatic, and acetic acids. The solutions have an intense green colour. The sulphate crystallizes readily, in long needles or square plates. Prussiate of potash, when dropt into these solutions, occasions a milk white precipitate, hydrosulphuret of potash, a black precipitate, infusion of nut-galls no precipitate.
11. Cobalt, like nickel, forms two oxides, but the protoxide only is found as a constituent of salts. Its colour is blue, but its carbonate has a shade of red. It dissolves in sulphuric, nitric, muriatic, and acetic acids. The solution has a red colour, provided there be no excess of acid; but the muriatic solution, when there is an excess of acid, is of a deep green colour. Prussiate of potash, when dropt into these solutions, occasions a brownish yellow precipitate, hydrosulphuret of potash, a black precipitate, and the infusion of nut-galls, a yellowish white precipitate.
12. The protoxide of manganese forms colourless salts, with acids. The peroxide combines with sulphuric acid, and forms a red solution, which does not crystallize. If the protoxide of manganese be precipitated from a solution by a caustic alkali, it falls down white; but, when the precipitate is exposed to the air, it gradually changes its colour and becomes black. When it is precipitated by an alkaline carbonate, it retains its white colour even when dried. The carbonate of manganese dissolves readily in sulphuric, nitric, muriatic, and acetic acids. The solution is colourless. It is precipitated white by prussiate of potash, white by hydrosulphuret of potash; while the infusion of nut-galls occasions no precipitate.
13. Cerium forms two oxides, the white and the brown, both of which unite with acids, and constitute salts. The salts of the protoxide are colourless, those of the peroxide yellow or orange. When cerium is precipitated from its colourless solutions by an alkaline carbonate, the precipitate is silvery white. Caustic alkalies likewise precipitate it white; and, if the white powder be heated to redness in an open platinum crucible, it becomes brown. Carbonate of cerium dissolves readily in sulphuric, nitric, muriatic, and acetic acids. The solutions have a sweet taste; they are precipitated white by prussiate of potash and hydrosulphuret of potash, by oxalate of ammonia, and tartrate of ammonia, while the infusion of nut-galls occasions no precipitate.
14. The protoxide of uranium seems alone capable of uniting with acids. The salts which it forms have a lemon yellow colour. Their solution in water is precipitated blood red by prussiate of potash, brownish yellow by hydrosulphuret of potash, and chocolate brown by infusion of nut-galls. The pure alkalies occasion a yellow precipitate, and the alkaline carbonates a white precipitate, soluble in an excess of the carbonate. The colour of the oxide, after being heated to redness, is greyish black.
15. The oxide of zinc forms with acids white salts, which are mostly soluble in water, and the solutions are colourless. These solutions have an astringent taste. Caustic potash occasions a white precipitate, which is redissolved by an excess of the alkali. An alkaline carbonate likewise occasions a white precipitate. A white precipitate is thrown down by prussiate of potash and hydrosulphuret of potash, while infusion of nut-galls occasions no precipitate.
16. Many of the salts of lead are insoluble in water. If a little of any of these be placed upon charcoal, and exposed to the action of the blow-pipe, a globule of lead is easily obtained from them. The soluble salts form colourless solutions, which have a very sweet and somewhat austere taste. These solutions are precipitated white by prussiate of potash, sulphate of soda, arseniate of potash, oxalate of ammonium, tartrate of potash, and infusion of nut-galls. They are precipitated black by hydrosulphuret of potash. The alkalies occasion a white precipitate, which is redissolved by adding an excess of alkali. The alkaline carbonates occasion a white precipitate, which becomes yellow when exposed to a red heat. 17. As few of the salts of tin can be obtained in permanent crystals, they are not likely to occur in the state of dry salts; though solutions of that metal in acid are not unlikely to occupy the attention of the practical chemist. Solutions of protoxide of tin are usually colourless; those of peroxide have often a shade of yellow. They are precipitated white by prussiate of potash; brownish black by hydrosulphuret of potash, if they contain protoxide; but golden yellow, if they contain peroxide of tin. Corrosive sublimate throws down a black precipitate from solutions of the protoxide, and a white from those containing the peroxide. Muriate of gold throws down a purple precipitate from solutions of the protoxide.
18. The salts of copper have a green or blue colour, unless they contain the protoxide of that metal, when their solution in water is colourless, and the crystals which they form are white. But, as the protosalts of copper are not permanent in the open air, they are not likely to occur to the experimenter except in a state of solution; and such solutions are easily recognised, because, if we expose a little of them to the open air, it soon acquires a green colour. Most of the salts of copper dissolve in water, or become soluble by the addition of nitric or muriatic acid. When ammonia is poured into the solution, it assumes a deep blue colour. Prussiate of potash occasions a red precipitate, hydrosulphuret of potash a black precipitate, and infusion of nut-galls a brown precipitate. When a plate of iron or zinc is plunged into a solution of a cuprous salt, the copper is gradually precipitated in the metallic state. Potash precipitates oxide of copper blue, and the precipitate is not redissolved by adding an excess of potash; but ammonia immediately redissolves it, forming a deep blue solution.
19. The salts of bismuth are colourless. They are insoluble in water; or at least when we attempt to dissolve them. The acid is separated by water, and the oxide left in the state of white flocks. The salts may be dissolved in acids. From these solutions prussiate of potash throws down a white precipitate with a shade of yellow, hydrosulphuret of potash a dark brown precipitate, and the infusion of nut-galls a light yellow precipitate.
20. Both the oxides of mercury combine with acids. Most of the mercurial salts are white; though some of them have a yellow colour. The solutions of all of them that dissolve in water are colourless. The insoluble salts are easily volatilized by heat in the state of a white smoke. Prussiate of potash, when dropt into a mercurial solution, occasions a whitish precipitate, which becomes yellow when exposed to the air; hydrosulphuret of potash occasions a black precipitate, and infusion of nut-galls an orange and yellow precipitate. Potash occasions a black or a brown precipitate, according to the state of oxidation of the mercury. When copper is put into a solution of any mercurial salt, the mercury is precipitated in the metallic state, and tinges the copper white.
21. Few of the salts of silver are soluble in water, or capable of being exhibited in the state of crystals. When exposed to the action of the blow-pipe on charcoal they are decomposed, and a globule of silver obtained. Those salts which are soluble in water are colourless, form transparent and colourless Decomposition, Chemical solutions, and are exceedingly acid. Prussiate of tin, Chropotash occasions in them a white precipitate, hydrosulphuret of potash a black precipitate, infusion of nut-galls a yellow precipitate, and common salt a white curdy precipitate, which becomes dark coloured when exposed to the light. When a plate of copper is put into a solution of silver, the metal is precipitated in the metallic state.
22. Few of the salts of gold have been hitherto formed. Hence they are not likely to come in the way of the experimenter. They have a yellow colour. Their solution in water is yellow. They are acid, and tinge the skin of a deep purple colour, which is indelible. When a plate of tin is put into a solution of gold, the liquid acquires a fine purple colour. When sulphate of iron is poured into it, the gold is precipitated in the metallic state.
23. The salts of platinum have scarcely been more studied than those of gold. Those of them that are soluble in water give the liquid a deep brown colour, and render it opaque. Several of the insoluble salts of platinum are yellow or brown, and when exposed to a red heat the platinum is reduced to the metallic state. Neither prussiate of potash nor infusion of nut-galls occasions any precipitate when dropt into solutions of platinum. Sulphured hydrogen occasions a black precipitate.
24. The salts of palladium, when dissolved in water have a fine red colour. Prussiate of potash occasions in these solutions a yellowish brown precipitate, hydrosulphuret of iron a blackish brown precipitate, and the alkalies an orange coloured precipitate. Mercury and the sulphate of iron throw down the palladium in the metallic state. When muriate of tin is dropt into a very diluted solution of palladium, it assumes an emerald green colour.
25. The solutions of the salts of rhodium in water are also red. No precipitate is occasioned in them by prussiate of potash, hydrosulphuret of potash, sal-ammoniac, or the alkaline carbonates. But the pure alkalies throw down a yellow powder, soluble in an excess of alkali.
26. The solutions of the salts of iridium in water are green; but when concentrated in an open vessel they become red. Prussiate of potash and infusion of nut-galls render the solutions colourless; but occasion no precipitation.
27. Tellurium, when in the state of an oxide, is capable both of combining with acids and with bases; so that it acts at once the part of a base and of an acid. The solution of these salts in water is colourless. No precipitate is formed in it by prussiate of potash; but hydrosulphuret of potash occasions a brown precipitate, and the infusion of nut-galls a flaky precipitate of a yellow colour. Zinc, iron, or antimony, precipitate the tellurium in the state of a black powder, which resumes its metallic brilliancy when rubbed.
28. The only salt of antimony at present known is tartar emetic, in which the tartar seems to act the part of an acid. Its solution in water is colourless. It is precipitated white by prussiate of potash, and orange-coloured by hydrosulphuret of potash. A plate of zinc or iron speedily throws down a black powder, especially if there be an excess of acid in the solution.
29. The salts of titanium form colourless solutions in water. The alkaline carbonates throw down from these solutions a white flaky precipitate. Prussiate of potash occasions a grass green precipitate. When an alkali is dropt in after the prussiate, the precipitate becomes purple, then blue, and at last white. Hydrosulphuret of potash occasions a dirty bottle-green precipitate. Infusion of nut-galls occasions a bulky blood-red precipitate. When a rod of tin is plunged into a solution of titanium, the liquid around it gradually assumes a fine red colour. But a rod of zinc occasions a deep blue colour.
Order II.—Combustible Salts.
These salts may be divided into the following genera, deriving their names from the acids of which they are partly composed.
1. Acetates. 8. Tartrates. 2. Pyrotartrates. 9. Citrates. 3. Formates. 10. Saclactates. 4. Zumates. 11. Urates. 5. Benzoates. 12. Malates. 6. Succinates. 13. Sorbates. 7. Oxalates. 14. Lactates.
Let us see how these different genera are to be distinguished from each other.
1. An acetate is easily recognised by putting a little of it into sulphuric acid, and applying a moderate heat. The fumes of acetic acid are driven off, and they make themselves known by their smell.
2. If a pyrotartrate be mixed with sulphuric acid, and distilled in a retort, an acid liquor comes over, and towards the end of the distillation a white sub-lime rises which is pyrotartaric acid. When the liquid in the receiver is exposed to spontaneous evaporation, crystals of pyrotartaric acid are formed in it. When this acid is dropt into acetate of lead or nitrate of silver no precipitate falls; but needle form crystals gradually make their appearance in the acetate of lead. It occasions a precipitate when dropt into the nitrate of mercury.
3. Formic acid agrees with acetic in forming soluble salts with all the bases. But it is destitute of the smell of acetic acid, and the salts which it forms with bases crystallize in different shapes. Thus acetate of copper crystallizes in four-sided pyramids, formate of copper in six-sided prisms. But the properties of the formates have been hitherto investigated so incompletely, that very precise characters of them cannot be laid down. They are not likely to occur to the practical chemist.
4. All the zumates are soluble in water. When sulphate of zinc is dropt into their concentrated solutions, a white precipitate falls. None of the other salts form precipitates with the zumates. When zumic acid is obtained in a separate state, it cannot be made to crystallize.
5. Most of the benzoates are soluble in water. When nitrate of mercury, or nitrate of tellurium, is dropt into a solution of a benzoate, a white precipitate falls; but persulphate of iron occasions an orange precipitate. If a benzoate be digested with sulphuric acid, the benzoic acid is separated and floats up on the surface. It may be separated, washed with water, and ascertained by its properties; particularly its silky lustre, its solubility in alcohol, its little solubility in water, and its volatility.
6. Many of the succinates are soluble in water. If an insoluble succinate be digested with sulphuric acid, the succinic acid is separated, and may be thrown down by diluting the acid with a sufficient quantity of water, and its properties investigated. If into a solution of a succinate, persulphate of iron, as neutral as possible, be dropt, a flesh-coloured precipitate falls. Muriate of barytes occasions a white precipitate, muriate of lime no precipitate.
7. Many of the oxalates are insoluble in water. When an oxalate is soluble, muriate of lime throws down a copious white precipitate from the solution. If we digest this precipitate long enough with sulphuric acid, sulphate of lime is formed, and oxalic acid disengaged, which may be crystallized and detected by its taste and its form. When an oxalate is insoluble in water, if we digest it for a sufficient time in a solution of carbonate of potash, part of the oxalic acid will combine with the potash. The potash liquid may now be saturated with nitric or muriatic acids, and then mixed with nitrate or muriate of lime. A white precipitate will fall, which being treated with sulphuric acid, as above described, will betray the oxalic acid which it contains.
8. If into the solution of a tartrate we let fall a few drops of potash ley, a granular precipitate falls, consisting of tartar, easily detected by its taste and little solubility in water. Muriate of lime does not occasion an immediate precipitate, when dropt into a tartrate. An insoluble tartrate may be partially decomposed by digestion with carbonate of potash. If the carbonate be saturated with muriatic acid, and then mixed with muriate of lime, and boiled, tartrate of lime falls, which may be decomposed by sulphuric acid, and the tartaric acid being obtained in crystals, is easily recognised by its properties.
9. Most of the citrates are soluble in water. These solutions are not precipitated by muriate of lime nor of strontian. But muriate of barytes, and muriate of lead, form white precipitates when dropt into concentrated solutions of the citrates.
10. The saclactates, when in solution, are precipitated by the salts of lime, barytes, and strontian; and by the nitrates of silver, mercury, and lead. Hitherto they have been so imperfectly examined, that more decided characters cannot be pointed out. The best mode of proceeding would be to separate the saclactic acid, which could be easily recognised by its properties.
11. The urates are nearly all insoluble in water. The best mode of examining them would be to separate the base by means of an acid. The uric acid in that case separates in the state of a white powder. It is characterized by its solubility in liquid potash, and by the fine pink colour which it acquires when heated with nitric acid on a watch glass, just when the whole of the acid is evaporated away.
12. Most of the malates are soluble in water, and Decomposition, Chemical.
The greater number of them refuse to crystallize. Malate of lime is soluble in water; but when the solution is mixed with alcohol, the malate of lime precipitates in the state of a coagulum. This property enables us to convert any malate into malate of lime. We may then separate the lime by means of sulphuric acid, and the malic acid, when obtained in a separate state, may be recognised by its properties. It throws down mercury, lead, and silver, from their respective solutions, in the state of a white powder.
13. If a sorbate be soluble in water, the best method of ascertaining the nature of its acid, is to precipitate it by means of sugar of lead, to wash the precipitate, to decompose it partly by sulphuric acid, and to throw down the remainder of the lead by a current of sulphureted hydrogen gas. The sorbic acid, thus disengaged, is a colourless liquid, destitute of smell, but having a strongly acid taste. It does not crystallize, and when evaporated to dryness absorbs moisture from the atmosphere. It precipitates salts of lead, lime, and barytes.
14. The lactates are all soluble in water. Hardly any of them crystallize, but when evaporated to dryness, have the appearance of gum or mucilage. To obtain the acid in a separate state, the simplest method seems to be to decompose the salt by means of sulphuric acid. Alcohol, digested on the dry mass, will dissolve the lactic acid, and leave the sulphate. Any little portion of sulphuric acid that may exist in the solution, may be removed by cautiously dropping in lactate of barytes. Lactic acid, when thus disengaged, has a brownish yellow colour, a sour taste, but no smell till it is heated, when it acquires a smell somewhat similar to that of sublimed oxalic acid. It is not precipitated from water by acetate of lime, nitrate of silver, muriate of lime, muriate of barytes, nor indeed by any salt whatever.
Sect. II. Of the mode of analyzing Salts.
After we have ascertained the genus and species of a salt, by attending to the characters laid down in the preceding section; the next step is to determine the proportions of the constituents of the salt. The analytical processes are greatly facilitated by the knowledge of several general principles, which have been lately ascertained, and which we shall therefore state in the first place.
1. All the acids and all the bases combine with each other in certain determinate proportions; so that a given weight of an acid being taken, we can assign the weight of each of the bases, which is just sufficient to neutralize it. These relative numbers are conceived to represent the weight of an atom or integrant particle of each acid or base. The following table represents these atoms, according to the best information which we at present possess.
1. Weight of the Atoms of the Acids.
| Acid | Weight of Atom | |-----------------------|---------------| | Carbonic acid | 2.75 | | Sulphurous acid | 4. | | Hyposulphurous | 3. | | Nitrous | 5.75 | | Chloric | 9.5 |
2. Weight of the Atoms of the Bases.
| Base | Weight of Atom | |-----------------------|---------------| | Ammonia | 2.125 | | Potash | 6. | | Soda | 4. | | Lime | 3.625 | | Barytes | 9.75 | | Strontian | 6.5 | | Magnesia | 2.5 | | Yttria | 5. | | Glucina | 3.25 | | Alumina | 2.125 | | Zirconia | 5.625 | | Protoxide of iron | 4.5 | | Peroxide of do. | 10. | | Protoxide of nickel | 4.375 | | Protoxide of cobalt | 4.625 | | Protoxide of manganese| 4.5 | | Protoxide of cerium | 6.75 | | Peroxide of cerium | 14.5 | | Protoxide of uranium | 16.625 | | Oxide of zinc | 4.125 | | Oxide of bismuth | 9.875 | | Protoxide of lead | 14. | | Protoxide of tin | 8.375 | | Peroxide of copper | 10. | | Protoxide of mercury | 26. | | Peroxide of mercury | 27. | | Oxide of silver | 14.75 | | Peroxide of gold | 27.875 | | Peroxide of platinum | 25.625 | The inspection of this table shows us that four parts by weight of sulphurous acid are neutralized by six parts by weight of potash, four parts of soda, five parts of yttria, ten parts of peroxide of copper, and so on. The weight attached to the acid is always neutralized by the weight attached to the base. Therefore, if a salt be neutral, and we ascertain the name of the acid and base by which it is composed, the inspection of the preceding tables will enable us to determine the proportions in which its constituents are united.
2. Many acids are capable of uniting in more than one proportion with bases. In such cases we always find, that one atom of base is united with one atom of acid, or with two atoms, or with three atoms, or with four atoms, or, in some rare cases, two atoms of base with three atoms of acid. In general, when there is more than one atom of acid united with one atom of base, the salt has the property of reddening vegetable blues, and it is distinguished by an acid taste. Such salts are called supersalts; and they are distinguished from the neutral salts by prefixing to the name of the neutral salt the syllables bi, tri, quadri, &c., indicating the number of atoms of acid united to one atom of base. Thus, bisulphate of potash is a compound of two atoms of sulphuric acid with one atom of potash; or, it consists, as we see from the preceding table, of 10 by weight of sulphuric acid united to six of potash. Quadroxalate of potash is a compound of four atoms of oxalic acid with one atom of potash; or it consists of eighteen by weight of oxalic acid united to six of potash; or, which is the same thing, it is a compound of three parts by weight of oxalic acid and one part of potash.
In such salts the number of atoms of acid and base united together, can only be determined by analysis; but the analysis is very much facilitated by the knowledge of the general law.
3. Some bases are capable of uniting in more than one proportion with acids. In such salts the same observations apply as in the supersalts. One atom of acid unites with one, two, or more atoms of base. Such salts, if they happen to be soluble in water, which is not always the case, have the property of giving a green colour to vegetable blues, and exhibit other alkaline properties. These combinations are distinguished by the name of subsalts. They are named precisely in the same way as the supersalts; but the syllable sub is always prefixed to show, that it is the number of atoms of base which is increased, not of the acid. Thus subbi-borate of soda is a compound of one atom of boracic acid and two atoms of soda. A much smaller number of subsalts is known at present than of supersalts: whether this be owing to a smaller number actually existing, or to the attention of chemists not having been so much turned to the subsalts as to the supersalts, it would be premature in the present state of our knowledge to determine.
4. Besides the acid and base, many salts contain also a certain proportion of water, usually distinguished among chemists by the name of the water of crystallization. It was supposed at first, that no salt can crystallize, unless it contains water as a constituent. But this opinion is not well founded. For many crystallizable salts contain no water whatever. This is the case with nitre, with common salt, and with sulphate of potash. Few of the salts of lead contain any water; yet many of them crystallize with facility, or are found crystallized in the earth. The quantity of water of crystallization varies prodigiously in different salts, and the proportions do not seem reducible under any general law. Some salts, which contain much water of crystallization, are deliquescent, as muriate of lime; others are efflorescent, as sulphate of soda; while others are not altered by exposure to the atmosphere, as alum. The water of crystallization exists always in the same proportion in the same salt; but the number of atoms vary in different salts, from 1 atom, which is the minimum, to 36, or even a greater number of atoms.
5. Sometimes two salts have the property of uniting together in a certain determinate proportion, and of forming a new salt, differing in its properties from both its constituents. Such compounds are usually called triple salts; though the term double salts would be more appropriate. Thus alum is a compound of 3 atoms of sulphate of alumine, and 1 atom of sulphate of potash, together with 23 atoms of water; and salt of Seignette, or Rochelle salt, is a compound of 1 atom of tartrate of potash, and 1 atom of tartrate of soda. Almost all the salts of ammonia and magnesia, with the same acid, are capable of uniting together, and of forming double salts.
Let us now consider the method of subjecting the salts to a rigid analysis. As the greater number of salts contain an acid, a base, and a quantity of water, we must, in order to solve the problem with accuracy, ascertain the proportion of all these three constituents.
1. If the salt is capable of bearing a red heat, without separating its acid from its base, we can readily ascertain the proportion of water of crystallization, by heating a given weight of the salt in a platinum crucible, and determining the loss of weight. Let it be required, for example, to determine the water of crystallization of sulphate of soda. Put into one scale of an accurate balance a platinum crucible, and balance it exactly, by putting the requisite weights in the other scale. Now, put 50 grains of the dry crystals of sulphate of soda into the crucible, and 50 grains in weights in the opposite scale, so as still to maintain the balance. Sulphate of soda contains so much water of crystallization, that when the crystals are heated they melt. Consequently, if the platinum crucible, containing the 50 grains of salt, were suddenly exposed to a strong heat, the salt would boil, and part of it would be driven out of the crucible, which would make the water of crystallization appear greater than the truth. To prevent this loss, the crucible is to be placed upon a Decomposition, Chemical.
Steam bath, or upon a sand bath, not heated above the boiling temperature, and kept in that position, till so much of the water of crystallization is dissipated, that the whole salt is reduced to a solid state. The crucible is then put into a fire, covered with a lid, and the heat raised till the salt is thoroughly red hot, and in a state of fusion. The crucible is now to be withdrawn from the fire, and allowed to cool. When quite cold, it is to be well wiped, and placed again in the balance. It will now be lighter than the weights which formerly counterpoised it in the opposite scale. Add weights till the counterpoise is again restored. These weights will indicate the water of crystallization of the salt. If the experiments be rightly conducted, and the salt in the requisite state, the water of crystallization, in 50 grains of sulphate of soda, will be found to amount to 28 grains, which is equivalent to 56 per cent.
2. If the salt is not capable of bearing a red heat without decomposition, it is still possible, in many cases, to determine the water of crystallization. Suppose, for example, that the acid can bear a red heat without volatilization, but that the base is ammonia, which is usually driven off by a very moderate heat. Let the salt, for instance, be phosphate of ammonia, or tungstate of ammonia. Weigh out 50 grains of it, and put it into a small retort. This retort must be fitted to a very small receiver, containing a saturated solution of caustic potash. Before the experiment begins, the retort, and likewise the receiver, must be accurately weighed, and their weights noted down. Heat is then gradually applied to the bottom of the retort, till the salt becomes visibly red hot. The whole is then weighed a second time. The loss of weight sustained by the retort is equivalent to the weight of the base of the salt, and to that of the water of crystallization which the salt contained. The receiver will have increased in weight, and the increase will be just equal to the water of crystallization of the salt. This increase subtracted from the loss of weight sustained by the retort, will give the weight of ammonia which the salt contained.
Suppose 50 grains of acetate of potash to be presented to us for examination, and that it is required to determine the water of crystallization which this salt contains. The potash, in the present case, is capable of bearing a red heat, but not the acetic acid. We may proceed in this way: Mix the 50 grains of acetate of potash with an excess of sulphuric acid in a platinum crucible, and expose the mixture at first to a low heat, which is to be very slowly raised till the crucible is heated intensely red hot. The heat of a common fire will not be found sufficient to dissipate the excess of acid, unless it be urged by a pair of bellows. By this process the acetate of potash is changed into sulphate of potash. Ascertain the weight of this new salt, 6/7ths of it are potash. Thus we obtain the weight of the potash contained in the 50 grains of acetate of potash. Let this weight be \(a\). Now, the weight of an atom of acetic acid is 6.375, that of potash is 6. Hence, we have \(6 : 6.375 :: a : \text{quantity of acetic acid in 50 grains of acetate of potash} = x\). From this proportion we obtain \(x = \frac{6.375a}{6}\). Now, it is obvious, that \(a + x\) is the amount of the potash and acetic acid in 50 grains of acetate of potash. If we subtract this sum from 50, the remainder will give us the water of crystallization contained in the salt.
Many of these salts, though incapable of bearing a red heat without decomposition, may be exposed to the heat of boiling water without injury. And in many cases, this temperature is sufficient to remove the whole or nearly the whole of the water of crystallization. Then in many cases this method may be applied with success.
4. When both of the constituents of a salt are volatile, and when it cannot be exposed even to a moderate heat without volatilization, in such a case we are unable at present to determine the water of crystallization directly. Such is the case with acetate of ammonia. But even in this case we are not without resource; though the only methods which can be followed require uncommon skill and every possible precaution. If we dissolve the acetate of ammonia in water, and add potash to the solution (noting the quantity) till we have just driven off the whole of the ammonia, and saturated the acetic acid; knowing the weight of potash employed, we can hence determine the weight of the acetic acid saturated. Knowing the weight of acetic acid, we can deduce from the preceding table the weight of ammonia requisite to saturate it. The weight of the acetic acid being added to that of the ammonia, and the sum subtracted from that of the original quantity of acetate of ammonia experimented on, the remainder will be the weight of the water of crystallization which that salt contains.
II. The second step in the analysis is to determine the weight of acid contained in a determinate weight of the salt. Three different modes are pursued, according to the nature of the acid. We must give an account of each.
1. When the acid is such that it forms an insoluble precipitate with some particular base. In that case we precipitate the acid by means of a salt containing the requisite base, wash and dry the precipitate, and ascertain its weight. This (as the composition of the precipitate is known) enables us to determine the weight of acid which the salt contained.
To give an example: Sulphuric acid is completely precipitated from every soluble salt by mixing the solution of the salt with muriate of barytes. Sulphate of barytes is a compound of five sulphuric acid and 9.75 barytes, so that \(\frac{5}{14.75} = 0.339\) of the precipitate, supposing it well washed and dried in a red heat, is sulphuric acid. Let it be required to determine how much sulphuric acid exists in fifty grains of crystallized sulphate of soda. Put the fifty grains of the salt into a cylindrical glass vessel, about twelve inches in length and four inches in diameter. Pour over the salt twenty times its weight of distilled water, and stir the liquid occasionally with a glass rod, till the salt is completely dissolved. Have in a clean phial a solution of muriate of barytes in distilled water. Pour a little of this liquid into the solution of the sulphate of soda, and then stir the solution with a glass rod. It will become immediately milky. Continue to add muriate of barytes as long as the milkiness appears to increase. Then leave the mixture for an hour or two till it has deposited the whole of the sulphate of barytes formed, and till the liquor has become clear. Let a drop of muriate of barytes fall into it; if it does not become white, we may conclude that the whole sulphuric acid has been precipitated. If it becomes white, we must add an additional quantity of muriate of barytes, and stir the liquor again with a glass rod. When it has become clear again, we try it once more with muriate of barytes. In short, muriate of barytes must be added till the supernatant liquor is not affected by any farther addition of that salt. When this is accomplished, allow the glass cylinder to remain at rest till the whole of the precipitate is deposited, and the liquid is perfectly transparent. Then draw off the clear liquor, either by means of a small syphon, or by the little glass instrument represented in fig. 16. It consists of a glass hollow ball, terminating both above and below in a glass tube. The undermost part of the lower tube is drawn out into a capillary fineness. The method of using this little instrument, usually called a sucker, is to plunge the capillary end into the liquid, and, applying the mouth to the other extremity, to suck in the liquor by the lips till the ball is filled with it. Then applying the tongue rapidly to shut the end in the mouth, the sucker is lifted out, and its extremity being put into another glass vessel, the tongue is withdrawn, upon which the liquid again runs out. This may be repeated till the whole colourless liquid is removed. Considerable promptitude is requisite in applying the tongue. For if any of the liquid be allowed to run out before the sucker be withdrawn, it will agitate the precipitate, and, rendering the whole liquor muddy, will prevent you from proceeding with the process till it has had time to settle. Another form of this little instrument is represented in fig. 17. But it is only when salts are analyzed upon a smaller scale than fifty grains, that the sucker can be employed with advantage.
When the clear liquid has been thus withdrawn as completely as possible from the precipitate, without running the hazard of losing any portion of the powder, a quantity of distilled water is to be poured into the glass, and the precipitate is to be stirred with a glass rod till it is well mixed with the water. A paper filter, formed by folding a sheet of blotting paper into the quarto form, and then with a pair of scissors cutting off the corner farthest from the centre into the segment of a circle, so as to give the folded paper the appearance represented in fig. 18.; it is obvious that if this paper were unfolded, it would have a circular shape. There is a paper prepared on purpose for filtering, and sold by apothecaries under the name of filtering-paper. But nothing answers so well as printing paper, procured from the mills before it has been sized. The filter is to be put in a glass funnel, placed upon a wooden stand, with a vessel below it to receive the liquid. A plan and elevation of such a stand, with funnels in it, is represented in fig. 19. The number of funnels in the stand may be increased at pleasure. In the figure, three only are represented. Four is the usual number.
The liquid containing the precipitate is to be poured upon the filter, and care must be taken to wash out the whole precipitate from the cylindrical glass with distilled water, and to pour the whole upon the filter, without allowing a single drop to escape. When the whole water has run through the filter, an additional quantity of distilled water is to be poured upon the precipitate. This must be continued till the water which passes through the filter is perfectly tasteless, and till it ceases to render muddy a solution of sulphate of soda, when a drop of it is let fall into that liquid. When this is the case, it is a proof that the sulphate of barytes on the filter is sufficiently washed. When the filter ceases to drop, it is to be removed from the funnel and placed upon some folds of blotting paper. The corners cut off from the paper, in preparing it for a filter, will answer the purpose. It may now be dried in the open air, or what is still better, by putting it upon a steam-bath, which is a large square plate of iron heated by steam. When dry, it is to be laid upon a table till it has become quite cold, and till it has imbibed the quantity of moisture which the air of the room supplies. The filter then, with the sulphate of barytes which it contains, is to be exactly weighed in an accurate balance. We then, by means of a blunt platinum spatula, or ivory spatula, loosen the precipitate from the paper. We then put as much of this precipitate as we can into a platinum crucible of known weight, and we ascertain the weight of sulphate of barytes thus introduced. Let the weight be $a$. The platinum crucible is now to be put into a fire, covered with its lid, and exposed for half an hour to a good red heat. When cold we weigh it again. The loss of weight indicates the quantity of water driven off from the sulphate of barytes by the red heat. Let the weight of the sulphate of barytes thus dried be $b$. Take the dry filter and rub off the whole of the remains of the precipitate with a dry cloth, then weigh the filter again. The weight of the filter subtracted from the original weight when the filter contained the sulphate of barytes, gives us the whole weight of the sulphate of barytes before it was exposed to a red heat. Let this weight be $c$. Let the weight of the whole sulphate of barytes, after it has been exposed to a red heat (supposing it possible to make the experiment), be $x$; then we have $\frac{a}{b} = \frac{c}{x}$, and, of course, $x = \frac{bc}{a}$.
In the case of sulphate of barytes, it is possible to separate the whole of the precipitate from the filter. But, in many cases, this cannot be done. In such cases, we remove as much of the precipitate as we can; we then measure the size of the filter, and burn it to ashes in a platinum crucible, and ascertain the weight of these ashes. A filter made from the same paper, and of the same size, is likewise to be burnt to ashes, and the weight of these ashes ascer- This last quantity being subtracted from the first, gives us the quantity of precipitate that had remained attached to the filter.
When the precipitate left on the filter has the property of absorbing moisture, as is often the case, we must weigh the filter while still warm. But this cannot be done in the usual way, because the filter would be constantly absorbing moisture and becoming heavier. To prevent this, procure a glass tube shut at one end, and having a cork covered with sealing wax fitted to the open end. This tube must be of such a size as to contain the filter when properly wrapped up. The weight of the tube and its cork should be written upon it with a diamond. In such a tube the filter may be weighed conveniently.
Having thus ascertained the weight of the sulphate of barytes, we have only to multiply that weight by 0.339; the product will be the quantity of sulphuric acid contained in the sulphate of soda under analysis.
There is no accurate method of determining, by analysis, the weight of the soda contained in the sulphate of soda; but when we know the weight of water and of sulphuric acid which the salt contains, we may infer the weight of the soda; since that weight, added to the two others, will constitute the sum total of the weight of the salt analyzed. Thus, suppose we find 50 grains of sulphate of soda to yield
| Water | 28 gr. | |-------|--------| | Sulphuric acid | 12.2 | | Total | 40.2 |
It is obvious that 9.8 grains are wanting to make up the total weight 50; hence we infer, that the soda in 50 grains of sulphate of soda weighs 9.8 grains.
We may take another example of an analysis of a salt, conducted in exactly the same way; and we shall select common salt, as very convenient for our purpose.
Common salt contains no water of crystallization; but it is seldom destitute of water mechanically mixed with it; therefore, before subjecting it to analysis, it is proper to expose it for some time to a red heat in a platinum crucible; then weigh out 50 grains of it, and dissolve them in twenty times their weight of water.
The solution is to be mixed with a solution of nitrate of silver in distilled water, as long as any precipitate falls. The precipitate, when well washed and dried, is a compound of chlorine and silver, in the proportion of 4.5 chlorine to 13.75 silver; so that
\[ \frac{4.5}{18.25} \text{ of the precipitate, or } \frac{0.2465}{18.25} \text{ of it is chlorine.} \]
Hence the weight of the precipitate, multiplied by 0.2465, gives us the weight of chlorine in 50 grams of common salt; this weight, subtracted from 50, leaves the weight of the sodium, which is the other constituent of common salt.
Phosphoric acid may be completely precipitated from the phosphates, by muriate of lime. When the precipitate is properly washed and dried,
\[ \frac{100}{180.5} \text{ or } 0.554 \text{ of its weight is phosphoric acid.} \]
The fluoric acid is likewise precipitated by muriate of lime; but the quantity of fluoric acid indicated by a given weight of precipitate, has not been ascertained with the requisite degree of accuracy.
2. When the acid is of such a nature that it can be completely precipitated from the base with which it is combined, by means of some other acid, and thus obtained in a separate state, the simplest method is, to precipitate it, wash it, and dry it. It is easy then to ascertain its weight, and thus to determine the proportion of acid which a given weight of the salt subjected to analysis contains. In this way may the acids in all the salts, whose acid is insoluble in water, or nearly so, be ascertained. Thus the acid of the borates may be precipitated by sulphuric acid; molybdic and tungstic acids may be precipitated by nitric acid; arsenious acid, in many cases, may be precipitated by sulphuric acid or muriatic acid; uric acid may be precipitated by acetic acid; sialactic acid may be precipitated by muriatic acid.
3. When the acid is of such a nature that it may be destroyed by heat, while the base remains unaltered, the method is to deprive the salt, in the first place, as completely as possible, of its water of crystallization. A given weight of it is then put into a platinum crucible, and exposed to a heat, at first low, but gradually raised to such a degree of intensity as to enable it to dissipate the acid completely, and to leave nothing behind but the base. Suppose it was required to determine the quantity of acid in a given weight of oxalate of lime. The oxalate is first made as dry as possible by exposing it for several hours to a temperature equal to that of boiling water. It may be afterwards left for 24 hours under the exhausted receiver of an air-pump, in which there is placed a wide shallow vessel, containing sulphuric acid. The salt thus dried is put into a platinum crucible, and exposed at first to a red heat. The temperature is then raised to whiteness, to dissipate the carbonic acid formed and united to the lime in the first part of the process. The residual lime being now weighed, and its weight subtracted from the original weight of the dried salt, the remainder indicates the weight of oxalic acid which was originally united to the lime.
In this way may combinations of the vegetable and animal acids, with various acids, be analyzed. When the acid is united to a metallic oxide, easily reducible to the metallic state, as the protoxide of lead, in order to obtain accurate results, it is necessary to dissolve the residual oxide in acetic acid. That portion of lead which has been reduced to the metallic state during the process will remain unacted upon. This quantity must be weighed, and \( \frac{1}{3} \) of its weight must be added to it. This additional weight being united to the weight of oxide left, after the decomposition of the salt, will constitute the true weight of the oxide, which being subtracted from the original weight of the salt, will give the quantity of acid with which the oxide of lead was originally united in the salt.
III. The last step in the process is the determination of the weight of base with which the acid is united. In almost all the salts, except those of potash, soda, and ammonia, the base may be thrown down by means of these alkalies, or their carbonates. And the precipitate being washed and thoroughly dried, indicates the weight of base which the salt contained. In many cases it is more correct to employ an alkaline carbonate as a precipitant than a pure alkali. In this way lime, barytes, and strontian, are very conveniently thrown down from their combination with acids. The precipitates thus obtained consist not of the pure earths, but of the earths united to carbonic acid. So that a portion of the weight of the precipitate must be subtracted before the true weight of the base is known. In carbonate of lime the carbonic acid amounts to $\frac{275}{637.5}$ or 0.431 of the whole; in carbonate of barytes it is $\frac{275}{1250}$ or 0.22; in carbonate of strontian to $\frac{275}{925}$ or 0.297. Hence if we multiply the weights of the respective precipitates by these numbers, the products obtained, subtracted from the original weight, will give us the true weight of lime, barytes, or strontian, contained in the salt analysed.
In some cases the oxide which constitutes the base of the salt, changes its state after its precipitation. Thus the protoxide of iron thrown down from any of the proto-salts of iron, cannot well be dried without being converted into peroxide of iron. Now, when 4.5 of protoxide of iron is converted into peroxide, its weight is increased to 5. Let the weight of peroxide obtained be $a$, and that of the protoxide required be $x$. We have $\frac{5}{4.5} = \frac{a}{x}$. Hence $x = \frac{4.5}{5}a$; or it is nine-tenths of the peroxide.
We have already, in a preceding part of this chapter, pointed out the method of determining the quantity of ammonia, or of potash, or soda, which exist in a salt.
**CHAP. III.—Of Mineral Waters.**
The analysis of waters consists in determining the different substances, whether saline or gaseous, which waters hold in solution. The processes are minute and tedious, and require very great precision. We shall divide this chapter into two sections. In the first section we shall point out the method of determining the quantity and nature of the gaseous contents of mineral waters; in the second section we shall point out the mode of detecting the saline contents, and of ascertaining the amount of each.
**Sect. I.—Of the Gases contained in Waters.**
Suppose a mineral water is presented for analysis, and we are required to determine whether it holds any gaseous matter in solution; and if it does, to ascertain the nature and proportions of the gaseous contents. It is proper to observe, in the first place, that the gases held in solution in waters, are of so fugitive a nature, that we cannot expect to detect them in any water, unless we examine it fresh from the source. If the water has been brought from a distance, though in well stopped bottles, we may be pretty certain that the gaseous constituents, or at least a considerable part of them, have made their escape during the transit of the water. If the water be fresh drawn, the method of proceeding is this: Have a retort, the capacity of which is known. Fill it with the water in question. Plunge its beak under a jar filled with mercury on the mercurial trough, and then by means of a lamp cause the liquid in the retort to boil as long as any gas is extricated from it. The gas, as it is extricated, passes into the glass jar along with the vapour of the water; and after the process is at an end, its quantity may be easily estimated, as the jar into which it has passed must be supposed divided into cubic inches and tenths. The bulk of the water being known, and that of the gas extricated being known, we have the proportion of gas contained in the water under examination. Thus, suppose that the experiment is made upon 25 cubic inches of water, and that the gaseous product measures 2.5 cubic inches: we infer, that every 100 cubic inches of the water contains 10 cubic inches of gas.
The only gases hitherto observed in mineral waters (at least capable of being separated by the above method) are,
1. Oxygen. 2. Azote. 3. Carbonic acid. 4. Sulphureted hydrogen.
Those waters that contain sulphureted hydrogen are easily distinguished by the smell of that gas which they exhale, by their property of blackening silver, and of striking a black when mixed with acetate of lead.
The waters which contain carbonic acid gas have an acidulous taste, a sparkling appearance, like champagne or bottled small beer. When dropt into lime water they render that liquid milky. If we attempt to dissolve acetate of lead in them, the solution becomes milky; but the addition of a little acetic acid renders them perfectly transparent.
For the mode of determining the quantities and species of these gases when extricated, we refer to the first chapter of this article.
**Sect. II.—Of the Saline contents of Waters.**
To determine the proportion of salts contained in a mineral water, the simplest, easiest, and most accurate mode, is to put a determinate weight of the water (1000 grains, or 10,000 grains, according to circumstances) into a wide and flat platinum, or silver capsule; to place the capsule upon a steam-bath, so regulated that the heat shall never exceed 180°. In this position the aqueous part will gradually evaporate away insensibly, without ever boiling, so that there will be no risk of any of the liquid being thrown out of the vessel. The saline contents will at last remain, forming a very thin crust upon the surface of the capsule. The capsule may now be removed to a sand-bath, and exposed to a heat sufficiently high to drive off all the water of crystallization. The weight of the platinum capsule being already known, we have only to weigh it carefully. When the salt is attached to it the increase of weight Decomposition indicates the quantity of saline matter furnished by the water under examination.
Upon the saline contents thus separated, pour a quantity of distilled water, equal to that in which the salts were originally dissolved. If the whole saline matter be dissolved in this water, the probability is that the saline matter has not been altered by its crystallization. But if a portion remain undissolved, as is usually the case, then we may conclude, as Dr Murray has very ingeniously suggested, that some of the salts have mutually decomposed each other, when brought into a concentrated state by the evaporation, and that salts have been formed which did not exist in the mineral water before its evaporation.
Almost the only salts combined in water are the carbonates, sulphates, and muriates of soda, lime, magnesia, and iron. The first point is to determine the different acids and bases present; the second to ascertain the relative weight of each.
Carbonic acid is ascertained by the water containing the carbonate forming a precipitate with muriate of barytes, which precipitate redissolves in nitric acid with effervescence.
When sulphuric acid is present, muriate of barytes occasions a precipitate not redissolved by nitric acid.
When muriatic acid is present, nitrate of silver produces a white curdy precipitate, which becomes dark-coloured when exposed to the light, and which is dissolved by means of caustic ammonia.
Lime is detected by the white precipitate produced in the water by oxalate of ammonia. Magnesia is precipitated by ammonia, or by dropping into the water, first carbonate of ammonia, and then phosphoric acid. The presence of iron is detected by the water containing it striking a black with infusion of nut-galls, and a blue with prussiate of potash.
The formula suggested by Dr Murray (Transact. Royal Society of Edinburgh, VIII. 259) is fully as accurate a means of analysing waters as any other, and seems to be easier of execution. It has been followed for many years by the writer of this article; but with one or two modifications, which seem calculated to guard against mistake. Supposing the gaseous ingredients of a mineral water to be ascertained, and the weight of its saline ingredients determined, we may proceed to the accurate analysis of it, as follows:
Measure out a determinate volume of it (as 50 or 100 cubic inches), and evaporate it, in an open vessel, down to one third. Divide this evaporated liquid into three equal portions. Precipitate the first portion by muriate of barytes; wash the precipitate, collect it, dry it at a red heat, and weigh it; digest it in nitric acid, dry it, and weigh it again. The loss of weight indicates the quantity of carbonate of barytes which the precipitate contained. The residual weight is sulphate of barytes; the carbonic acid in the water is equivalent to 0.22 of the weight of the carbonate of barytes; the sulphuric acid to 0.339 of the weight of the sulphate of barytes.
Precipitate the second portion by means of nitrate of silver; wash the precipitate, dry it, and fuse it in a platinum capsule, previously weighed. By weighing the capsule, containing the fused chloride of silver, the weight of the precipitate may be ascertained.
The fourth part of this weight is equivalent to the weight of the muriatic acid, contained in the portion of water precipitated.
Precipitate the third portion by oxalate of ammonia; wash and dry the precipitate; expose it to a red heat, in a platinum crucible; pour on it some dilute sulphuric acid; digest for some time, then evaporate to dryness, and expose the crucible to a pretty strong heat; weigh the sulphate of lime thus formed, 0.453 of its weight, indicates the quantity of lime in the portion of water precipitated.
Pour into the same third portion of the water thus freed from its lime, a portion of carbonate of ammonia, and let phosphoric acid fall into it, drop by drop, as long as any precipitate continues to fall. This precipitate being washed, dried, and exposed to a red heat, is phosphate of magnesia. 0.357 of the weight of this salt is equivalent to the weight of the magnesia contained in the water.
If the water be a chalybeate, a portion of it equal to one of the three preceding portions, must be taken and mixed with a solution of benzoate of ammonia. The precipitate being washed, dried, exposed to a red heat, and weighed, nine-tenths of its weight indicates the weight of protoxide of iron contained in the water.
By the preceding steps, the weight of all the substances contained in the water will be ascertained, except the soda. To know the amount of it, we cannot do better than follow exactly Dr Murray's formula. Evaporate a portion of the water to one-third. Precipitate the carbonic and sulphuric acids by muriate of barytes, taking care not to add any excess of the precipitant. Throw down the lime by oxalate of ammonia, and the magnesia by means of carbonate of ammonia and phosphoric acid. Then evaporate the liquid thus treated to dryness. A quantity of common salt will remain, which must be exposed to a red heat; 0.4 of this weight indicates the sodium contained in the bulk of water employed. And 0.4 sodium is equivalent to 0.53 of soda.
It seems hardly requisite to mention some other substances that occasionally make their appearance in mineral waters. Silica is a common constituent, very easily obtained in a separate state. We have only to evaporate a portion of the water to dryness, and redissolve the saline residue in distilled water. The silica will remain undissolved, and will betray itself by its insolubility in acids, and its easy fusibility into a transparent glass, with soda, before the blow-pipe.
CHAP. IV.—Of Metals.
Nothing is more likely to occur to the young analyst for investigation than alloys of various metals in various proportions. Before such alloys can be successfully analysed, it is necessary to be acquainted with the properties of all the metals in a pure state, because methods of separating them from each other can only be deduced from a knowledge of these properties. It will be requisite, therefore, to divide this chapter into two sections. In the first section, we shall point out the method of ascertaining the name of any pure metal presented for examination: in the second, we shall give the formulas for analysing the most important alloys.
Sect. I.—Method of ascertaining the Names of Metals.
The metals at present known amount to 33; all of which might be easily recognized, by determining their specific gravity, and observing their colour and hardness. The following table exhibits these properties, as far as they are at present known:
| Colour | Sp. Gravity | Hardness | |--------|-------------|----------| | Potassium | White | 0.865 | 4 | | Sodium | White | 0.972 | 4 | | Calcium | White | — | — | | Barium | White | — | — | | Strontium | White | — | — | | Magnesium | White | — | — | | Manganese | Grey | 8.013 | 8 | | Iron | Grey | 7.8 | 9 | | Zinc | White | 7.1908 | 6.5 | | Copper | Red | 8.895 | 7.5 | | Nickel | White | 8.82 | 8.5 | | Cobalt | Grey | 8.7 | 6 | | Uranium | Grey | 9 | 8 | | Palladium | White | 12.148 | 9 | | Tin | White | 7.299 | 6 | | Antimony | White | 6.712 | 6.5 | | Molybdenum | White | 8.611 | — | | Mercury | White | 13.568 | — | | Arsenic | White | 5.763 | 5 | | Tellurium | White | 6.115 | — | | Bismuth | White | 9.822 | 7 | | Lead | Blue | 11.352 | 5.5 | | Silver | White | 10.510 | 7 | | Tungsten | White | 17.4 | 9 | | Chromium | White | 5.9 | 9 | | Titanium | Yellow | — | — | | Columbium | Grey | 5.61 | 8 | | Cerium | White | — | — | | Osmium | Grey | — | — | | Gold | Yellow | 19.361 | 6.5 | | Platinum | White | 21.5313 | 8 | | Rhodium | White | 10.694 | 9 | | Iridium | White | 18.58 | 9 |
But as it may not be always in the power of the experimenter to ascertain these properties with accuracy, we shall point out the method of determining each metallic body from its chemical properties. Suppose a metal presented for investigation, and that it is required to determine its name: 1. The first thing to be done is, to put a portion of it into pure water. If an effervescence take place, and the metal dissolve in the liquid, we may conclude that it is an alkaline metal; or one of those in the first division of the preceding table. Now, as none of the alkaline metals, except potassium and sodium, has been hitherto obtained in a permanent state, we may conclude that our metal is one or other of these two. Saturate the alkali thus formed, by the solution of the metal in water, with sulphuric acid, and evaporate the liquid till it deposits crystals. It is very easy to distinguish whether these crystals be sulphate of potash, or sulphate of soda. Sulphate of potash forms very hard small crystals, which are fragments of the pyramidal dodecahedron. They require a great deal of water to dissolve them; and are not altered by exposure to the atmosphere. Sulphate of soda crystallizes in six sided prisms, terminated by dihedral summits, and the sides of the prism are channeled longitudinally. When these crystals are exposed to the air, they speedily effloresce, and fall to powder.
The metal called manganese likewise effervesces in water. But instead of dissolving, it is converted into a green-coloured oxide. The process is much slower with this metal than with potassium or sodium. But as manganese has not been applied to any useful purpose, and as it speedily oxidizes in the air, there is but little chance of its coming in the way of the practical chemist, at least for analysis.
2. If the metal undergoes no sensible alteration in water, we may conclude, that it is neither manganese nor an alkaline metal; but some one of the others in the preceding table. The next thing to be done is, to put it into sulphuric acid, previously diluted with twice or thrice its bulk of water. If it effervesces and dissolves in this liquid, the metal is either iron or zinc. If it is not acted on sensibly by the acid, it is one or other of the remaining metals in the table.
The solution of iron in sulphuric acid is readily distinguished from the solution of zinc in the same acid. The zinc solution is colourless; that of iron light-green. Infusion of nut-galls strikes a black, when poured into the iron liquid; but it produces no change upon the liquor of zinc. Prussiate of potash throws down iron blue, but zinc white. When an alkali is dropped into the solution of iron, a greenish white precipitate appears in flocks, which soon becomes dark green, and finally changes to reddish yellow; but an alkali throws down the solution of zinc chalk white, and the colour undergoes no subsequent change.
3. Supposing the metal to undergo no change in dilute sulphuric acid, the next step of the investigation will be to put it into dilute nitric acid. If it effervesce with that acid, either cold, or when slightly heated, it will be one or other of the following fourteen metals:
1. Copper, nickel, cobalt, uranium, palladium. 2. Tin, antimony, molybdenum. 3. Mercury, arsenic, tellurium. 4. Bismuth. 5. Lead, silver.
We have divided these fourteen metals into five sets, each of which are characterized by certain peculiarities, which are easily recognized.
(1) The five metals, constituting the first set, when dissolved in nitric acid, tinge that liquid of some colour or other; while the solutions of all the others are colourless, like water. The solution of copper is blue, or greenish blue. When potash or soda is dropped into the liquid, a blue deposit falls, which is not redissolved by adding an excess of alkali. A few drops of ammonia occasion a bluish white precipitate, which is redissolved by adding an excess of the ammonia, and the liquid assumes a fine purplish blue colour. Prussiate of potash occasions a red precipitate, and a plate of iron plunged into the liquid is speedily covered with a coating of copper.
The solution of nickel is grass-green. When ammonia is poured into it in sufficient quantity, it becomes blue. Potash or soda throws down a green precipitate. Prussiate of potash occasions a milk-white precipitate. A plate of iron when put into it does not throw down any metal.
The solution of cobalt is violet red. The alkalies throw down from it a blue-coloured precipitate. When this precipitate is mixed with borax, and fused before the blow-pipe, it forms a violet blue glass. The prussiate of potash occasions a brownish yellow precipitate. Infusion of nut-galls produces a yellowish white precipitate. A plate of iron throws down no metal.
The solution of uranium is yellow, and when concentrated by evaporation, it yields lemon yellow crystals. The alkalies throw down a yellow precipitate from this solution; prussiate of potash, a blood red precipitate; and infusion of nut-galls, a chocolate-coloured precipitate. A plate of iron does not throw down any metal.
The solution of palladium is red. When protosulphate of iron is dropped into it, the palladium is precipitated in the metallic state. Prussiate of potash throws down an olive; the alkalies an orange precipitate. Muriate of tin renders the solution opaque, by throwing down a brown precipitate; but if the liquid be sufficiently diluted, it assumes a fine emerald green colour.
(2.) The three metals, tin, antimony, and molybdenum, which constitute the second set in the preceding series, have this peculiarity, that when treated with concentrated nitric acid, they are not dissolved, but merely converted into white or yellowish white powders. We have it not, therefore, in our power to examine the liquids when these metals are the subjects of experiment. But the characters of tin, antimony, and molybdenum, are so different, that they are very readily distinguished from each other.
Tin is a ductile metal, and it is characterized by a creaking noise which it emits when bent, or when squeezed between the teeth. It dissolves with effervescence in muriatic acid. When this recent solution is mixed with muriate of gold, a beautiful purple powder precipitates, well known by the name of purple of Cassius. Corrosive sublimate, dropped into the recent solution of tin in muriatic acid, occasions a black precipitate; but when dropped into the permuriate it throws down a white precipitate. Prussiate of potash produces a white precipitate. Hydrosulphuret of potash occasions a brownish black precipitate when dropped into the protomuriate of tin; a golden yellow when dropped into the permuriate. Antimony is a very brittle metal, and easily reduced to a fine powder in a mortar. It dissolves with great rapidity in nitro-muriatic acid. The solution has a yellowish colour. When water is poured into it, the liquid becomes milky, and a copious white precipitate falls down. Prussiate of potash, when dropped into the nitro-muriatic solution of antimony, occasions a white precipitate, while hydrosulphuret of potash occasions an orange-coloured precipitate. When a plate of iron is put into the liquid, a black precipitate gradually falls.
Molybdenum is a brittle metal; but so infusible, that it has been hitherto obtained only in grains. The white powder into which it is converted by nitrid acid, possesses the properties of an acid. When diffused through water, it reddens vegetable blues, and combines very readily with the alkalies. If a plate of zinc or of tin be put into water through which this powder is diffused, the powder gradually assumes a blue colour.
(3.) The three metals, mercury, arsenic, and tellurium, form colourless solutions in diluted nitric acid; but they are easily distinguished from the other three metals, that likewise form colourless solutions, by their great volatility.
Mercury being always in a fluid state at the common temperature of the atmosphere in this country, and being the only known fluid metal, cannot be confounded with any of the other metallic bodies, even by the most inexperienced experimenter. Hence it would be needless to point out its chemical properties when in solution.
The properties of arsenic are scarcely less striking than those of mercury. It is a very soft and brittle metal. If we put it into a retort and expose it to heat, it is completely volatilized, though the temperature does not exceed 356°. It is deposited in crystals in the throat of the retort. When laid upon a red hot iron in the open air, it burns with a pale blue flame, and evaporates in a white smoke, having the smell of garlic.
Tellurium is at present so scarce a metal, that it is not likely to come in the way of a chemical analyst, at least in the metallic state. It is very brittle and volatile. When heated before the blow-pipe it burns with a blue flame, and evaporates in a white smoke, having the smell of radishes. When an alkali is dropped into the solution of tellurium, in nitric acid, a white precipitate appears, which is redissolved on adding an excess of alkali. Prussiate of potash occasions no precipitate, infusion of nut galls a yellow precipitate, and hydrosulphuret of potash a brownish black precipitate. A plate of iron throws down the tellurium in the state of a black powder, which resumes the metallic lustre when rubbed.
(4.) Bismuth, the only metal belonging to the fourth set of the above series, dissolves very rapidly in nitric acid of moderate strength. The solution usually crystallizes in the course of twenty-four hours. If the solution of bismuth in nitrid acid be diluted with a great deal of water, it becomes milky, and a copious precipitate of a white powder, which is a hydrate of bismuth, falls.
(5.) The fifth set in the preceding series contains only lead and silver, two metals so different from each other, and so familiarly known to every one, that it seems scarcely necessary to point out the chemical differences between them.
The solution of lead has a sweet taste, accompanied with a certain degree of astringence. When evaporated sufficiently, it crystallizes in octahedrons or tetrahedrons of a white colour, opaque, but having considerable lustre. When these crystals are exposed to a red heat a yellow oxide remains, which readily melts when heated, and flows into a glass of a dark-reddish yellow colour, and a considerable degree of hardness. Nitrate of lead is precipitated white by sulphuric acid and the sulphates, white by prussiate of potash, white by infusion of nut-galls, and black by hydrosulphuret of potash.
The solution of silver, on the contrary, in nitric acid, has an acrid taste, and constitutes one of the most corrosive substances known. When sufficiently concentrated by evaporation, it crystallizes in transparent plates, which melt easily when heated, without undergoing decomposition. The solution of nitrate of silver is precipitated in white floccs, like curd, by common salt. This precipitate becomes speedily dark-coloured when exposed to the light. It is dissolved by ammonia. The solution of protosulphate of iron, poured into nitrate of silver, precipitates the silver in the metallic state.
4. If the metal resists the action of diluted nitric acid, even when assisted by heat, the next step in the investigation will be to expose it to a red heat, in contact with air, in a platinum crucible. If by this process it is converted into an oxide, we may conclude it to be one or other of the following six metals:
- Tungsten, chromium, titanium, columbium, cerium, osmium.
When tungsten is calcined in the open air, it is converted into a yellow-coloured powder, which possesses the properties of an acid, and is known by the name of tungstic acid. If we mix tungsten with its weight of nitre, and calcine the mixture in a crucible till the nitric acid is decomposed, the residual mass is almost completely soluble in water. Muriatic acid, when dropt into this solution, throws down a white precipitate, which, when boiled in an excess of acid, becomes yellow, and possesses the characters of molybdic acid.
Chromium, in the metallic state, is so scarce, that it is not likely to come in the way of the experimenter. By calcination it is converted into a green powder, scarcely soluble in acids. If chromium be mixed with its own weight of nitre, and kept for half an hour at a red heat in a crucible, a yellowish-coloured mass is obtained, which dissolves in water, and the solution has a yellow colour. Nitrate of silver, dropt into this liquid, throws down a purple precipitate; nitrate of lead a brilliant orange precipitate; and nitrate of mercury a red precipitate.
Titanium has hitherto been scarcely ever obtained in the metallic state, so that it cannot well offer itself for the investigation of the practical chemist; it is hardly necessary, therefore, to point out its characters. When calcined in contact with air, it assumes a blue colour. It dissolves in nitro-muriatic acid by the assistance of heat. The solution, when freed as much as possible of all excess of acid, has a pale yellow colour. Infusion of nut-galls throws down from it a blood red precipitate. Prussiate of potash throws down a grass green precipitate mixed with brown. When an alkali is dropt in after the prussiate, the precipitate becomes purple, then blue, and, at last, white. When a rod of tin is plunged into the nitro-muriate of titanium, the liquid around it gradually assumes a fine red colour; a rod of zinc, on the other hand, occasions a deep blue colour.
Columbium has been reduced to the metallic state by Professor Berzelius. When calcined it undergoes an incipient combustion, and is converted into a greyish white matter; but it cannot, by this method, be completely oxidized. When mixed with nitre, and thrown into a red hot crucible, a feeble detonation takes place. A snow-white mass is thus obtained, which must be digested in muriatic acid, to separate the potash. The white substance remaining after this treatment is columbic acid. When fused with eight times its weight of carbonate of potash, it forms a compound which dissolves in water. Muriatic acid throws down the columbic acid in the state of a white hydrate. Neither prussiate of potash nor hydrosulphuret of potash occasion any precipitate; but infusion of nut-galls throws it down orange, provided there be no excess either of acid or alkali in the liquid.
Cerium having not hitherto been obtained in the metallic state, except in very minute portions, cannot come under the examination of the experimental chemist. It dissolves in nitro-muriatic acid, when assisted by heat. The solution, when freed as much as possible of its excess of acid, has a sweet taste. It is precipitated white from this solution by prussiate of potash, hydrosulphuret of potash, tartrate of potash, oxalate of ammonia, and the alkalies. Infusion of nut-galls occasions no precipitate. When the precipitate by oxalate of ammonia, or tartrate of potash, is calcined in a platinum crucible, it assumes a reddish brown colour.
When osmium is heated in the open air, the oxide evaporates as it forms, exhaling a very strong odour, and acting very powerfully upon the eyes. If osmium be mixed with its own weight of nitre, and heated in a retort, a white sublimate rises, having a strong odour, a caustic taste, and soluble in water. The solution is colourless, but becomes blue when mixed with the infusion of nut-galls. When a rod of zinc, or a quantity of alcohol or ether, is poured into the solution, the osmium precipitates in white floccs.
5. If the metal under examination is neither altered by water, sulphuric acid, nitric acid, or calcination in the open air, it must be one or other of the following:
- Gold, platinum, rhodium, iridium.
Gold is easily known by its yellow colour and its great specific gravity. It dissolves with great facility in nitro-muriatic acid. The solution has a yellow colour, a caustic taste, and tinges the skin purple. Protomuriate of tin precipitates it purple. Protosulphate of iron precipitates the gold in the metallic state. When ammonia is poured into the solution a yellow precipitate falls in floccs, which, when dried, assumes the form of a brown powder, and fulminates loudly when exposed to heat.
Platinum has a white colour and a higher specific gravity than any other metal. It dissolves in nitro-muriatic acid, provided the acids of which the liquid is composed be as concentrated as possible. The solution is dark brown and opaque. Sal ammoniac throws down a fine yellow precipitate from this liquid. When the precipitate is exposed to a red heat, it is converted into metallic platinum. The nitro-muriate of platinum is neither precipitated by prussiate of potash, nor infusion of nut-galls; but sulphureted hydrogen throws down a black precipitate.
Rhodium is a white brittle metal, which, when pure, is not dissolved by nitro-muriatic acid, nor by any acid compound whatever. It may be obtained dissolved in acids, when the solution is facilitated by the presence of certain other metals. The solution, in such cases, is red, and it is not affected by prussiate of potash, hydrosulphuret of potash, or infusion of nut-galls. The pure alkalies throw down a yellow powder, soluble in excess of alkali.
Iridium dissolves in concentrated nitro-muriatic acid, but with difficulty. When ammonia is poured into the concentrated solution, a great number of small brilliant crystals are deposited, of so deep a purple, that they appear black, and capable of communicating an orange red colour to a great quantity of water. This colour is destroyed by dropping into the liquid protosulphate of iron, or sulphureted hydrogen, or by putting into it a rod of iron, zinc, or tin.
Sect. II.—Method of Analysing some of the principal Alloys.
It would be needless to lay down formulas for the analysis of all the conceivable alloys of the metals, as the greater number of them could never be of any practical utility. We shall satisfy ourselves with noticing such of them as are applied to purposes of utility, and which, therefore, must frequently offer themselves to the view of the practical chemist.
1. Silver Coins.
The modern European silver coins are all alloys of silver and copper, in very different proportions, according to the country where they are issued. Some idea of the proportion of copper in these coins may be formed from the following table, drawn up by Dr Thomson, from a very extensive series of analyses of these coins. (Nicholson's Journal, XIV. 409.)
| Alloy per cent. | Weight of Silver, that of the Copper being 1. | |----------------|---------------------------------------------| | British | 7.5 | | Dutch | 8 | | French | 9 | | Austrian | 9.5 | | Sardinian | 9.5 | | Spanish | 10.5 | | Portuguese | 11 | | Danish | 12 | | Swiss | 21 | | Russian | 24 | | Hamburgh | 50 |
Besides silver and copper, it is not uncommon also to find traces of gold in silver coin.
To analyse a silver coin, we may put it (previously washed with soap and water, and well wiped, to remove any grease with which it may be covered) into about four times its weight of strong and pure nitric acid, previously diluted with three or four times its weight of water. It dissolves rapidly, with effervescence. The gold remains undissolved in the state of a fine powder. It may be collected on a watch glass, well washed, carefully dried, and its weight ascertained. It will never be found to exceed a minute fraction of a grain.
The nitric acid solution contains the silver and the copper, which last metal gives it a blue colour. The silver is to be precipitated by means of common salt, and the precipitate being well washed, dried, and exposed to a heat sufficient to fuse it, is to be weighed with the requisite accuracy. $\frac{13}{17.5}$ or $0.753$ of this weight is silver.
Into the liquid thus deprived of its silver, pour a quantity of sulphuric acid, sufficient to decompose the nitrate of copper which it contains. Evaporate the liquid nearly to dryness, and dissolve the whole residue in distilled water. Put a rod of zinc into the liquid, and allow it to remain for 24 hours. The whole of the copper will be precipitated in the metallic state. Wash the zinc rod, to separate any copper which may adhere to it. Allow the copper to settle to the bottom of the vessel. Decant off the colourless liquid, and replace it by distilled water, acidulated with muriatic acid. An effervescence will take place, owing to the solution of a portion of zinc which had fallen down along with the copper, and there is usually, at the same time, a quantity of nitrous gas exhaled. The copper may now be collected upon a watch glass, well washed, dried, and weighed.
If this analysis has been rightly conducted, the weights of the gold, silver, and copper, added together, ought to equal the weight of the original coin subjected to experiment.
In some of the silver coins of the Roman emperors, as of Alexander Severus, and Gordianus, there exists a little lead along with the preceding constituents. When such coins are dissolved in nitric acid, the oxide of the lead remains mixed with the gold. The gold may be separated by digesting the residue in nitro-muriatic acid. We may then boil the residual oxide of lead in muriatic acid, and evaporate to dryness. The muriate of lead thus formed being dried, heated, and weighed, $\frac{13}{17.5}$ or $0.743$ of its weight will be lead.
In the copper coins of Gallienus, we find a quantity of tin. When the coin is dissolved in nitric acid, the oxide of tin remains undissolved. It may be dissolved in muriatic acid, precipitated in the metallic state by a rod of zinc, and weighed.
When copper coins contain lead as well as tin, which is the case with the Chinese coins, the process of analysis is the same, only the lead must be precipitated from the nitric solution by means of sulphate of soda. The white precipitate being well washed... What is called pewter in this country was perhaps originally pure tin, but it has long been customary to alloy the tin with some other metal. Lead constitutes by far the most common alloy; but antimony is said, also, sometimes to exist in pewter. But, whether pewter be an alloy of tin and lead, or of tin and antimony, the process of analysis is nearly the same. Put the pewter, to be analysed, into a phial containing three or four times its weight of concentrated nitric acid, and allow the liquid to remain in contact with the alloy, at a boiling temperature, till all action on it has ceased. Let us suppose the pewter, in the first place, to be an alloy of tin and lead.
When the nitric acid has ceased to act upon the pewter, we must put the whole into a porcelain evaporating dish, and evaporate nearly to dryness. Distilled water is then poured upon the white mass, the whole is digested for some time, and then thrown upon a filter. The white powder which remains on the filter is a perhydrate of tin. It must be washed with water till that liquid passes off tasteless, then dried, and exposed to a red heat. It thus acquires a yellow colour, and is peroxide of tin \( \frac{7.375}{9.375} \) or \( \frac{0.786}{1} \) of the weight of this oxide indicates the tin in the portion of pewter subjected to analysis.
The watery solution contains the nitrate of lead. It must be all collected, and concentrated by evaporation. Being mixed with sulphate of soda, the lead is thrown down in the state of sulphate of lead. This sulphate, being edulcorated, washed, and exposed to a red heat, and weighed, gives the weight of the lead contained in the pewter. For \( \frac{13}{19} \) or \( \frac{0.684}{1} \) of dry sulphate of lead is lead.
When the pewter consists of tin and antimony, nitric acid converts it into a white hydrate, without dissolving any of it. But, if we wash away the nitric acid, and substitute muriatic acid in its place, the whole white matter is dissolved. If water be poured into the solution, the antimony is precipitated in the state of a white hydrate. Collect this hydrate, dry it, and expose it to a red heat. It is now deutoxide of antimony. \( \frac{100}{123.7} \) or \( \frac{0.808}{1} \) of this oxide is equivalent to the weight of the antimony in the pewter. The tin may be precipitated from the muriatic acid in the metallic state by a rod of zinc, and its weight ascertained.
3. Brass.
Brass is a compound of copper and zinc. There are two varieties of it. The first, or British brass, composed of an atom of copper, and an atom of zinc. The second, or Dutch brass, is composed of two atoms of copper, and 1 atom of zinc. The mode of analyzing both varieties is the same.
Dissolve the brass in diluted nitric acid, and pour into the solution a considerable excess of liquid potash. The copper is precipitated in the state of a blue hydrate, while the zinc is retained in solution by the potash. Expose the hydrate of copper to a red heat, and weigh it; four-fifths, or \( \frac{0.8}{1} \) of its weight is equivalent to the weight of the copper contained in the brass.
Pour a slight excess of sulphuric acid into the liquid containing the zinc in solution. Then add carbonate of potash, as long as any precipitate falls. The white precipitate, which is carbonate of zinc, is to be well washed, dried, and exposed to a red heat. By this exposure, it is converted into oxide of zinc. \( \frac{4.125}{5.125} \) or \( \frac{0.8048}{1} \) of this oxide is equivalent to the weight of the zinc contained in the brass.
If the brass happen to contain lead, we can detect the presence of that metal by dropping sulphate of soda into the nitric acid solution. The sulphate of lead will precipitate, and we can deduce the quantity of lead from it by the rule already laid down.
If the brass contains iron, its presence will be indicated by dropping prussiate of potash into the nitric acid solution; a blue precipitate will fall, or at least the liquid will assume a blue tinge. The peroxide of iron may be precipitated by the cautious addition of ammonia to the nitric acid solution, previously rendered neutral by evaporation to dryness, and solution in distilled water. This peroxide being washed, dried, exposed to a red heat, and weighed, \( \frac{3.6}{5} \) or \( \frac{0.7}{1} \) of its weight is the iron which existed in the brass.
4. Bronze and Cannon Metal.
Bronze is an alloy of copper and tin. This alloy was used by the ancients for making into cutting instruments. It is often called brass in English. The mode of analysis is nearly similar to those already described.
The bronze is to be dissolved in concentrated nitric acid, and the action of the acid is to be promoted towards the end of the process, by raising it to the boiling temperature. The copper will dissolve, while the tin will remain in the state of peroxide. The weight of the two metals thus separated from each other may be determined by the rules already laid down in this section.
The analysis of bell-metal is to be conducted exactly in the same way as that of bronze. But bell-metal, besides copper and tin, often contains zinc, and sometimes silver.
When bell-metal is treated with nitric acid, the tin remains in the state of a perhydrate. The copper, zinc, and silver are dissolved. Precipitate the silver by means of common salt, and determine its quantity by the rules above laid down. The nitric acid solution now contains only copper and zinc, which may be obtained in a separate state by the method pointed out for analysing brass.
5. Silver and Lead.
Lead, when first obtained from its ores, is almost always alloyed with a certain quantity of silver. This metal renders the lead much harder and less ductile. It is, therefore, always removed, when the quantity of silver contained in the lead is sufficient to defray the expense of the process. The proportion of silver is always determined by the process called cupellation. As this process is employed likewise to ascertain the purity of silver and gold, and to obtain silver in a state of purity for the purposes of the silversmith, it will be worth while to describe it in this place. It may be performed either in a common reverberatory furnace, or in the muffle of a wind furnace, heated to the temperature of about 270° Wedgewood.
Cupels are small, flat, porous cups, weighing about 200 grains, and composed of bone ashes, well calcined, pounded, and washed. This powder is mixed with a proportion of fern ashes, and is moistened and beat with mallets into an iron mould. If a metallic alloy be placed into such a cupel, composed of two metals, one of which is easily oxidized by heat and air, while the other undergoes no alteration, and if the oxide be easily fusible into a liquid glass, this liquid oxide will make its way through the pores of the cupel, while the metal, not having any affinity for the cupel, will remain behind in the state of a metallic button.
Place the cupel in the furnace, and bring it to the requisite degree of heat; then place upon it a certain determinate quantity of the lead to be subjected to cupellation. The lead speedily melts; becomes covered with a coat of oxide; exhales fumes, and undergoes a considerable agitation, which, by renewing the surface, promotes the process of oxidation. In this way the whole of the lead is gradually converted into an oxide; this oxide melts, and is absorbed by the cupel. As the oxidation proceeds, the bulk of the lead diminishes, and it leaves a reddish brown trace upon the surface of the cupel. The surface, which was at first flat, becomes evidently convex, and brilliant points may be seen on it, which are continually increasing in size. At last, when the lead is nearly quite removed, the brilliant points disappear, the metallic button becomes covered with an iridescent pellicle, which disappears instantaneously, and the brilliant metallic lustre is again restored. The process is now finished. But care must be taken not to cool the silver button too suddenly, lest a portion of it should be squeezed out and lost by the too rapid congelation of the external coating. The button of silver is now to be weighed, and its weight indicates the proportion of silver which the lead under examination contains.
Silver may be freed from copper, or any other metal with which it is alloyed (except gold and platinum), by the same process. The quantity of lead employed must increase with the proportion of alloy which the silver contains; varying from about twenty times to four times the weight of the silver, according to its purity or impurity. Indeed, the proportions of lead may be much smaller than are usually indicated in books, provided the process be conducted with the requisite skill. Place the lead upon the cupel, and when it is melted, wrap up the silver in a piece of paper, take it up with a pair of forceps, and place it over the middle of the lead. The two metals incorporate almost instantaneously. The phenomena of the cupellation are precisely similar to what has just been stated. The oxide of Decomposition, Chemical, lead carries along with it the whole of the copper, and leaves the silver in a state of considerable purity.
6. Gold and Silver.
Gold and silver are separated from each other by means of nitric acid. The alloy is laminated and put into hot nitric acid. The silver dissolves, but the gold remains behind in the state of a powder, which is easily fused into a button and weighed. It is to be observed, however, that it is not possible by this method to separate the whole of the silver from the gold, unless the weight of the first metal considerably exceeds that of the second. When this is not the case, we must, in the first place, fuse the alloy with such a quantity of silver as will make the amount of the silver in the alloy three times that of the gold. We then laminate this alloy, and treat it with nitric acid. Knowing the portion of silver added to the alloy, it is easy to determine its composition.
When gold is alloyed both with silver and copper, as is often the case with gold coins, the first process followed is to remove the copper by the cupel. The silver is then separated from the gold by means of nitric acid, adding previously the requisite proportion of silver.
We might now give the processes for separating the metallic oxides, when they happen to be mixed or united together in various proportions. But the methods already given for the metals will apply likewise to the oxides. It is necessary, however, to observe, that the greater number of oxides lose their power of dissolving in acids, if they have been exposed to a red heat. To restore their solubility, we must mix them with two or three times their weight of potash, or carbonate of potash, and expose them to a red heat. The potash may then be washed away with water, and the oxides will remain in a state of solubility in acids. Or if the potash has the property of dissolving the metallic oxide, we may precipitate the oxide by saturating the potash with an acid, and then pouring an alkaline carbonate to the solution, or by passing at once a current of carbonic acid gas through the potash solution.
Chap. V.—Of Stones.
The term stone is applied to the indurated masses of which the surface of the earth is composed. These masses have been carefully examined by mineralogists, and have been found composed of about 164 different species of minerals, to which names have been assigned. These minerals, with a very few exceptions indeed, may be considered as salts, often chemically united with each other, or mechanically mixed in various proportions. As we are not in possession of any method of distinguishing the chemical combinations from the mechanical mixtures, and do not even know any means of ascertaining, whether a mineral subjected to analysis be pure or impure, much less progress has been made in ascertaining the chemical nature of stones, than might have been expected, from the great labour and abilities bestowed on the investigation. The subject, therefore, still claims the closest attention of the practical chemist, and on that account it will be proper to give a pretty minute account of the best method of analysing these bodies with accuracy.
The acids found in stones are chiefly silica, sulphuric acid, carbonic acid, phosphoric acid, fluoric acid, boracic acid, arsenic acid, and tungstic acid. The bases are alumina, lime, barytes, strontian, magnesia, and, in some rare cases, yttria, glaucia, zirconia. Oxide of iron is a common constituent of stones. In some rare instances oxides of manganese, copper, chromium, and nickel, are found in them.
We shall divide this chapter into two sections. In the first section we shall give an account of the apparatus requisite for analysing stones; and, in the second, lay down the formulas necessary to be followed in conducting the analysis. Considerable practice is necessary to enable us to conduct the different processes without loss, or with the least possible loss. We should endeavour always to make the number of processes as few as possible; for the risk of loss and of error obviously increases with the number of processes through which our mineral has to go.
Sect. I.—Of the Apparatus for Analysing Stones.
Before a stone can be subjected to analysis, it is generally necessary to reduce it to an impalpable powder. We must, therefore, be provided with two mortars, which are appropriated to that purpose. The first of these is of steel, and is well known to jewellers by the name of the diamond mortar, because it is employed for pounding diamonds. It consists of three distinct pieces of polished steel, possessing a considerable degree of hardness. The first of these constitutes the bottom of the mortar. It is a circular piece of steel, about an inch thick. The top is flat and polished, only it has a circular ring raised a little above it, and the flat circular space within this ring is about an inch and a quarter in diameter. The second piece is a hollow cylinder of steel, of such a size as to fill exactly the circular space within the ring of the first piece, to which it is fitted by grinding. This cylinder is about two inches high, and the diameter within is about an inch. The third piece is a solid steel cylinder of the same length as the hollow cylinder, into which it is fitted by grinding, and it has a bulb on the top. The stone, previously broken into small pieces, is reduced to a pretty fine powder in this mortar. A little of it is put into the hollow cylinder, the solid cylinder is fitted in, and by a blow with a hammer it is smartly struck against the bottom part. By this blow the fragment is crushed, the pieces of the mortar are now removed, and the powder which lies on the bottom piece is carefully poured into a glass capsule. By proceeding in this way the stone is reduced to a coarse powder.
To convert this into an impalpable powder an agate mortar is employed. Plate LXIX. Fig. 20, exhibits an outline of this mortar. \(a\) is a section of it, \(b\) a profile of it, and \(c\) the pestle which is also of agate. These mortars are made at Oberstein in Germany, but may be purchased both in London and Paris, and cost half a guinea. They are about three inches in diameter, and about an inch and a half or two inches thick. These mortars are sufficiently hard for grinding in them the hardest stony bodies. But when the mineral to be pounded is very hard, as sapphire, spinel, &c., a portion of the mortar is always ground down at the same time. It is necessary to keep an account of this portion. This is done by weighing the mortar before and at the end of the process. The loss of weight which it has sustained, is the portion of the mortar which has been ground down and mixed with the mineral to be analysed. As agate mortars are composed almost entirely of silica, we have only to subtract from the weight of silica which we obtain from the mineral, the quantity which has been rubbed from the mortar. The remainder will be the silica really contained in the mineral. Suppose that, while pounding a mineral in our mortar, the mortar undergoes a loss of twenty grains, and that we obtain forty grains of silica by the subsequent analysis. It is obvious that our mineral will really contain only twenty grains of silica, the other twenty grains having been rubbed off the mortar. Mortars of rock crystal or flint might be likewise employed. But those of rock crystal would be more expensive than the agate mortars, without possessing any advantage over them: and flint is so very easily broken, even by a slight blow, that a mortar made of it would be apt to go to pieces in the very act of pounding.
To pound the mineral successfully in an agate mortar, some precautions must be taken. We must put only a small quantity (not more than a few grains in weight) into the mortar at once. The pounding is performed by rubbing the pestle against the bottom of the mortar; and we must continue the friction till all feeling of grittiness has disappeared, and till the powdered mineral adheres together in the form of a cake. As it is of the utmost consequence not to lose a particle of the mineral during the pounding, the mortar must be placed upon a sheet of clean paper, and we must be careful not to drive any of the mineral out of the mortar by incautious pounding. When the whole mineral is reduced to powder we must weigh it again. If the operation of pounding has been properly performed, the powder should be equal to the weight of the mineral and of the portion rubbed off the mortar both together.
When the mineral is reduced to an impalpable powder, the next step of the process is usually to mix it with twice or thrice its weight of caustic potash, or carbonate of soda, according to circumstances, and to expose the mixture to a red heat for an hour, in a metallic crucible. A crucible is a small vessel used by chemists and various artists, for exposing bodies to heat. It is made of clay, of black lead, of iron, silver, and platinum. Its shape is nearly conical, with the point of the cone cut off. The narrowest part constitutes the bottom and the widest the top. It is usually furnished with a lid composed of the same materials with the crucible itself. Its size varies from an inch or less to several feet in height, according to the purpose for which it is to be used. For the analysis of stony bodies, we should be provided with crucibles of fine silver and of platinum. The size may be about two inches in height, and an inch and a half in diameter at the mouth. Silver is the least acted upon by pure potash; but platinum is upon the whole preferable to silver, because it is not so easily fused, and because it is not injured, as silver is, when put into a common fire of pit-coal. For the sulphur which pit-coal usually contains, speedily combines with the silver and destroys it. When a silver crucible is to be used, it must be inclosed within a common clay or black-lead crucible, and care must be taken not to expose it to a heat sufficient to melt the silver. There is little risk of this in a common fire, if we do not urge it with bellows.
When the mixture of alkalis and stone has been exposed for an hour to a red heat, we must take it off the fire and allow it to cool. Upon inspecting this hot matter, we are enabled, from the appearance which it assumes, to draw some conclusions respecting the constitution of our mineral. If the mixture has melted completely, we may be sure that the mineral either consists entirely of silica, or at least contains a considerable proportion of that earth. If it has not melted, the probability is, that a great deal of alumina is present in the mineral. If the mixture while hot has a brownish red colour; but becomes green on cooling, we may be sure that iron is present. Grass green indicates manganese, and yellowish green chromium.
The next set of vessels required are dishes, into which the mixture is to be put, after being softened with water, in order to be dissolved in muriatic acid, and afterwards evaporated to dryness, or nearly so, upon a sand bath. The most convenient vessels for this purpose are those which are made of porcelain. In this country they are known by the name of Wedgwood's evaporating dishes. They are circular, and shaped somewhat like a saucer, only they have a spout at one side for the convenience of pouring out their contents, and they have no circular rim round the bottom. In the inside they are glazed, but destitute of glaze on the outside. The size varies from one inch in diameter to a foot or more. The most convenient, upon the whole, for mineral analyses, are about six inches in diameter, and about 2½ inches deep. The French chemists employ porcelain vessels, which have the shape of hemispheres, and are all glazed, both within and without, except the bottom, which is to be applied to the hot sand. These vessels are not so apt to crack as those of Wedgwood; but the glaze is more apt to crack in all directions. When this happens, it is impossible to make them completely clean; so that the same dish cannot be employed long for the analysis of minerals, without the risk of error from that cause.
The remainder of the apparatus employed in the analysis of minerals, does not require a particular description. The filters, the test-glasses, &c. are precisely similar to those described in a preceding part of this Article.
Sect. II.—Method of analysing Stones.
Let us suppose a stone to be presented for analysis. The first step of the process is to drop a fragment of it into muriatic acid. If it effervesces in the acid and dissolves, we may infer that it is an earthy Decomposition, Chemical. We shall first give the mode of analysing the more complicated stones. They are for the most part silicates, or combinations or mixtures of silicates.
I.—Simple Stones.
Great pains should be taken in selecting the specimen for analysis. It should always be as pure as possible, and should be selected for the purpose by a skilful mineralogist. The chemist, before he begins the analysis, should ascertain the specific gravity of the specimen, and write down a mineralogical description of it from actual observation. This serves as a kind of testimony of the nature of the mineral, after it has been destroyed by the analysis.
The quantity taken for analysis may be about 50 grains. When more than this is employed, the analysis becomes very tedious. The time is shortened in proportion to the smallness of the quantity examined. If we work upon a grain or two only, we may ascertain the nature of the constituents in an hour or two. But 50 grains enable us to determine the weight of each constituent with sufficient exactness, while it is so great as to render it unlikely that any one of the constituents should be overlooked, without our perceiving the oversight.
The first step of the analysis is the reduction of the mineral to an impalpable powder, in the way described in the last section. We then weigh out fifty grains of it, mix it with twice or thrice its weight of carbonate of soda, or pure potash, and expose it for an hour to a red heat in a platinum crucible. When the mineral contains a great proportion of silica, carbonate of soda answers just as well as caustic potash, while it is much cheaper. Klaproth, who first introduced the use of caustic potash in mineral analysis, always employed that alkali in solution in water, and gradually evaporated the mixture to dryness. He conceived that this rendered the mixture of the potash with the mineral more intimate, and of course facilitated the mutual action. In most cases litharge may be substituted for potash with advantage. It succeeds even when the mineral contains a considerable proportion of alumina. But we must be sure that the litharge is quite free from lead in the metallic state, otherwise it would injure the platinum crucible.
When the crucible is removed from the fire, and has become cold, we must wipe it quite clean on the outside with a cloth, and then, placing it in the middle of a Wedgwood evaporating dish, fill the crucible with distilled water. After standing some hours covered with the water, the mixture at the bottom of the crucible will be partly softened. Stir it up with a platinum spatula; pour off the water containing all the softened part of the mixture into the Wedgwood dish; fill up the crucible with distilled water again, and let it stand some hours as before. In this way you must proceed till the whole of the mixture has become soft, and has been washed out of the crucible.
Muriatic acid must now be poured into the watery liquid in the Wedgwood dish. An effervescence takes place, because the alkali has absorbed carbonic acid during the preceding process, which is driven off as it dissolves in the muriatic acid. You continue to add muriatic acid till all effervescence has ceased, and till the whole of the matter has dissolved.
Some minerals which contain little silica, and a great deal of alumina, will not be rendered quite soluble by heating them with an alkali. In such cases a portion of insoluble matter remains, upon which the muriatic acid has no action. This is a portion of the mineral in its original state. To render it soluble we must repeat the heating with an alkali, softening with water, and adding muriatic acid till we obtain a complete solution. Five or six repetitions of these processes may in some cases be requisite. With such minerals it is better to employ litharge, or borax, or phosphoric acid, than potash or soda. We will then obtain a complete solution by one process, which will greatly diminish the risk of error. If we employ too little alkali, it may happen that we do not obtain a complete solution in muriatic acid, even when our mineral contains a sufficient quantity of silica. The insoluble matter in such a case will be pure silica, and will readily enter into fusion when heated with a new portion of alkali.
Let us suppose the muriatic acid solution accomplished, the next step in the process is to evaporate this solution upon a sand-bath. Place the Wedgwood dish upon a sand-bath, and expose it to a heat sufficient to cause it to evaporate; but care must be taken not to raise the heat so high as to cause the liquid to boil, for in such a case a portion of it would be driven out of the dish and lost. When the liquid is considerably concentrated by evaporation, it loses its liquid form, and assumes that of a jelly, at least if it contains any considerable proportion of silica. As soon as this change has taken place we must stir it with a platinum spatula, and continue the agitation till the liquid is evaporated nearly to dryness. If this be neglected it is apt to sputter up, and part of it to be driven out of the vessel. Besides, the portion next the bottom is apt to be overheated, and some of the earthy salts might run some hazard of being decomposed. The evaporation need not be carried farther than the gelatinizing of the silica, if our sole object be merely to obtain the whole of that substance. For when silica is reduced to the state of a stiff jelly, it is no longer soluble in water.
Upon the gelatinous mass thus obtained distilled water is to be poured, the Wedgwood dish is to be put again upon the sand-bath, and the whole stirred about occasionally with the spatula, till the water has become almost boiling hot. The whole is now to be poured upon a filter. The silica having been rendered insoluble by the evaporation of the muriatic acid liquid to dryness, will remain upon the filter; but all the other constituents of the mineral, being in the state of muriates, will be dissolved in the water, and in that state will pass through the filter. Distilled water must now be poured upon the filter to wash the silica clean, and we must continue to do so till it passes through quite tasteless, and incapable of rendering a solution of common salt milky. The filter is now to be dried, and then accurately weighed. We then pour the silica into a platinum crucible, the weight of which has been previously noted. Wipe the filter quite clean with a cloth, and weigh it again. The difference of weight gives the quantity of silica collected on the filter. Let this weight be \(a\). By weighing the platinum crucible containing the silica, we find the quantity of it which we have collected in that vessel. Let it be \(b\). Expose the platinum crucible for half an hour to a red heat, and weigh it again as soon as it has become cold. The weight of the silica will be diminished; because it will, by the heat, be deprived of the whole water with which it was impregnated when weighed upon the filter. Let its new weight be \(c\). Let the weight of the whole silica, supposing it had been exposed to a red heat, be \(x\). We have
\[ \frac{b}{c} = \frac{a}{x} \]
Having thus obtained the silica in a separate state, the next step of the process is to separate and weigh the alumina. For this purpose we must pour carbonate of soda into the aqueous solution which has passed through the filter, till we throw down the whole of the constituents of the mineral that were united to the muriatic acid. This precipitate is to be edulcorated, and, while still moist, a quantity of caustic potash ley is to be poured over it, and the whole boiled for half an hour in a glass flask. The alumina, if any be present, will be dissolved in the ley, while the other substances will remain undissolved. Allow them to fall to the bottom. Decant off the potash ley; wash the residual powder clean with water, and add the liquid to the potash ley. Pour a solution of sal ammoniac into the potash till the liquid acquires a pretty strong smell of ammonia. The alumina will precipitate in white flocks. Continue to add sal ammoniac as long as any precipitate appears. The alumina thus separated must be washed, dried, and exposed to a red heat, and weighed.
Potash ley is capable of dissolving not merely alumina, but glucina also. The white powder thus precipitated by sal ammoniac, may therefore be glucina as well as alumina. To determine this point, we must dissolve it in sulphuric acid by the assistance of heat. If any portion remain undissolved, it must be considered as silica, and be added to the quantity of that substance found previously. Into the sulphuric acid solution we must pour a quantity of sulphate or muriate of potash, previously dissolved in water, and set the liquid aside for a few days. If the earth was alumina, we shall find a number of crystals of alum deposited at the bottom of the vessel. If it was glucina, no such deposit will have taken place. When all the crystals of alum that can be obtained have been deposited, we must wipe them dry and weigh them. The alumina which they contain, is equivalent to \(\frac{10.86}{100}\) or 0.1086 of the weight of the crystals. If this weight be equal to the whole of the earth originally dissolved Decomposition in the sulphuric acid, we may conclude that the whole of that earth was alumina. But if there be a deficiency, the probability is that some glucina was also present. This will be found in the sulphuric acid solution, from which it may be precipitated by an alkaline carbonate. We may be certain that this precipitate is glucina; if we find it soluble in liquid carbonate of ammonia; if it forms sweet-tasted salts with acids; and if it be precipitated white from acids by prussiate of potash, and yellow by infusion of nut-galls.
If yttria be suspected in the matter which remained undissolved after the action of the potash ley, we may digest it in carbonate of ammonia, which will dissolve that earth, but leaves all the other ingredients untouched. And the yttria may be obtained pure, and weighed, by evaporating the carbonate of ammonia to dryness, and exposing the residue to a red heat.
The residue of the mineral thus deprived of the silica, alumina, glucina, and yttria, may still consist of lime, magnesia, oxide of iron, oxide of manganese, and in some rare cases of oxide of chromium and oxide of nickel. It is always proper to determine, by means of tests, which of these bodies are present and which absent; because the mode of the analysis must vary with the number and nature of the ingredients. If, for example, nothing were present but lime and oxide of iron, we should dissolve the whole in muriatic acid, precipitate the iron by means of ammonia, and the lime by means of oxalate of ammonia. If the residue consisted of lime, magnesia, and oxide of iron, we should dissolve it in dilute sulphuric acid, evaporate the solution to dryness at as low a heat as possible; water, mixed with a little alcohol, would dissolve the sulphate of magnesia and iron, but would leave the sulphate of lime. This last sulphate, being heated to redness and weighed, would give us the lime contained in the mineral, which will amount to $\frac{3.625}{8.625}$ or 0.42 of the sulphates. The iron may be precipitated from the liquid solution by means of benzoate of ammonia. The precipitate being washed, dried, exposed to a red heat, and weighed, will be red oxide of iron. Nine-tenths of its weight are equivalent to the black oxide of iron which existed in the mineral. The magnesia may now be precipitated by potash and weighed.
If besides these three ingredients oxide of manganese be likewise present, the very same method of proceeding will answer; only, after having separated the iron by means of benzoate of ammonia, we must pour a little hydrosulphuret of potash or ammonia into the liquid which still contains the magnesia and manganese. By this addition the manganese will be precipitated. If it be heated to redness for some time in an open vessel, its weight will indicate peroxide of manganese. To convert it into protoxide of manganese we must multiply the weight of peroxide by $\frac{9}{11}$ or 0.818.
The chromium, when present, is indicated by protonitrate of mercury, forming a red precipitate. When this metal is suspected in the mineral subjected to analysis, from the peculiar green or red colour which it has, we must employ nitric acid instead of muriatic to form the original solution. After the silica is separated, and the liquid deprived of its excess of acid by the requisite evaporation, we may precipitate the chromium by means of protonitrate of mercury. The precipitate being dried, exposed to a red heat, and weighed, will give the quantity of protoxide of chromium. If the chromium was present in the state of chromic acid, to obtain its weight, we must multiply the weight of the green oxide by 1.3.
Nickel has hitherto been found in one stony mineral only; the chrysoprase, which consists chiefly of silica, and has an apple-green colour. In this mineral it is associated with iron and lime. After the separation of the silica and alumina, the iron may be precipitated by ammonia, and when it is separated, we may throw down the nickel by hydrosulphuret of potash. Nothing will remain but the lime, which may be thrown down by an alkaline carbonate or by oxalate of ammonia.
After having thus obtained all the constituents of the mineral in a separate state, and determined the weight of each, the next step is to add all these weights together. If they amount to the weight of the portion of mineral analysed, we have reason to conclude that the analysis has been rightly performed. But if there be a deficiency, we have either committed an error, or the mineral contains some ingredient which we have overlooked.
As water occurs very frequently in minerals, we must in the first place endeavour to discover whether the deficiency be not owing to a portion of that substance which we have not reckoned. For this purpose we must take a determinate weight of the mineral (fifty grains for example), and expose it for an hour to a strong red heat in a platinum crucible. The heat will drive off the water, if any be present, and the deficiency of weight, after the mineral has been allowed to cool, will indicate the quantity of water which has been driven off.
If the mineral contain no water, or if the quantity which it contains be insufficient to make up the deficiency between the original weight of the stone, and the weight of the constituents which we have obtained, it may probably contain an alkali. For three different alkalies have been found in stones; namely, potash, soda, and lithina. To detect this alkaline ingredient we must make a second analysis of the mineral, but we must conduct it in a different way. Fifty grains of the stone, reduced to a fine powder, must be mixed with four times its weight of nitrate of barytes, or three times its weight of carbonate of barytes. This mixture must be exposed for two hours to a strong red heat in a platinum crucible. If nitrate of barytes be employed, it will enter into fusion at a comparatively low heat. Of course, if the mixture were exposed suddenly to a strong red heat, it would swell greatly, and a portion of it would probably be driven out of the crucible and lost. To prevent this we must raise the heat gra-
CHEMICAL.
Dually, and not bring it to the utmost degree of intensity till the nitrate of barytes has had time to lose most of its acid.
The fused mass is to be taken from the fire, allowed to cool, softened with water, and dissolved in muriatic acid, precisely in the way described at the commencement of this section. Into the muriatic acid solution a quantity of sulphuric acid is to be poured, capable of decomposing all the muriates and converting them into sulphates. The liquid becomes immediately milky in consequence of the precipitation of the sulphate of barytes. Separate this precipitate by the filter. Then pour an excess of carbonate of ammonia into the liquid, and boil the whole for some minutes. All the earths and metallic oxides will be precipitated, and nothing will remain in the solution but the sulphate of ammonia, and the sulphate of the alkali contained in the mineral, if any such existed in it. Evaporate the liquid to dryness, and expose the dry mass to a red heat in a platinum crucible. The sulphate of ammonia will be sublimed, and nothing will remain but the alkaline sulphate derived from the mineral. Weigh this sulphate, then dissolve it in water, and crystallize the salt. It will be easy, from the shape of the crystals, and the properties of the salt, to determine whether it be sulphate of potash, sulphate of soda, or sulphate of lithina. The composition of all these sulphates being known, we may easily deduce from the weight of the sulphate previously ascertained, how much potash, soda, or lithina, our mineral contained.
If no alkali can be detected in the mineral, it may contain fluoric acid, which constitutes an ingredient of the topaz, and of some other analogous minerals. To ascertain whether any of this acid be present, mix a portion of the mineral reduced to fine powder with sulphuric acid, and expose the mixture to heat in a glass vessel. If the glass be corroded, and if the vessel acquires a smell similar to that of muriatic acid, we may conclude that fluoric acid is present. To determine its quantity we must make a new analysis of the mineral. Fifty grains of it are to be fused with an alkali softened with water, dissolved in muriatic acid, and the silica separated by the method described at the beginning of this section. The remaining liquid is precipitated by carbonate of potash, and being filtered, and exactly neutralized, is precipitated by means of lime water. The precipitate is fluor spar. It must be exposed to a red heat. 26.5 per cent. of its weight indicates the fluoric acid, if we consider fluor spar as a fluate of lime. But if we consider fluor spar to be a compound of fluorine and calcium, according to the hypothesis of Ampere and Davy, in that case the fluorine will amount to 46.69 per cent. of the fluor spar obtained.
II.—Carbonates.
As the carbonic acid cannot be conveniently collected and weighed, we are under the necessity of adopting a different method for the analysis of the carbonates. Provide a small crystal phial with two mouths, as in fig. 21. To one of those mouths let a crystal stopper be fitted. The other must remain open. Pour into this phial a quantity of concentrated nitric acid, recently heated to deprive it of the nitrous gas which the smoking acid of the shops always contains. Put into the mouth of the phial a plug of cotton wool. Balance this phial accurately upon the scales of a good beam. Suppose the carbonate to be subjected to analysis to be calcareous spar or common limestone. Break the mineral into small pieces, of such a size that they can conveniently pass through the mouth of the phial. But let there be no powder. Into the same scale that contains the phial with the nitric acid, put fifty grains of these pieces, and counterpoise them exactly by fifty grains weight put into the opposite scale.
The nitric acid must have been poured into the phial through the mouth furnished with a glass stopper. As soon as it is poured in, the mouth must be wiped with a piece of paper, and the stopper put in its place. Remove the cotton plug, and with a pair of forceps lift up the pieces of calcareous spar, and put them one after another into the phial through the open mouth. Then replace the cotton plug. The pieces will immediately begin to dissolve with effervescence, owing to the escape of the carbonic acid gas, and in proportion to its escape the weight will diminish, and the opposite scale will preponderate. When the solution is completed, or when the effervescence is at an end, remove the phial from the balance; place it upon a table; take out the glass stopper and the cotton plug, then introduce through one of the mouths of the phial a small glass tube, and plunge it nearly, but not quite, so low as the surface of the nitric acid. Apply the mouth to the other end of the tube, and blow air gently through it for about a minute. Then draw in air through it into the mouth for about another minute. This will remove the carbonic acid gas which is usually floating in the empty part of the phial, and materially affects the weight. Put the glass stopper and the cotton plug again in their places. Put the phial on the same scale of the balance where it was before, and add weights till the equilibrium is restored. These weights are equivalent to the weight of the carbonic acid which had made its escape during the solution of the mineral in the nitric acid.
By the same method may the quantity of carbonic acid present in carbonate of strontian, carbonate of barytes, carbonate of magnesia, carbonate of iron, magnesian limestone, and, indeed, in all the carbonates, be ascertained. When carbonate of barytes is analysed in this way, we must dilute the nitric acid with water, otherwise the solution does not succeed. This renders the result not quite so accurate as it would otherwise be; for when the nitric acid is very weak, it is capable of retaining a portion of the carbonic acid in solution. We may, indeed, determine the bulk of the quantity thus held in solution, by putting the liquid into a small flask, or retort, furnished with a bent tube, passing into a mercurial trough. By carefully heating the liquid, we can drive off the carbonic acid gas, and measure its bulk in a glass jar standing inverted over the mercury. But such an experiment must be made with great caution, lest we drive over nitric acid, which would act upon the mercury, and produce nitrous gas, the evolution of which would disturb all our estimates. The weight of carbonic acid in the carbonates being determined, we can ascertain that of the earthy bodies dissolved in the nitric acid, by the rules already laid down in a preceding part of this article. If there be any portion which refuses to dissolve in the nitric acid, we must carefully wash it, and dry it, and then heat it with thrice its weight of carbonate of soda. The fused mass may be analysed precisely in the way described at the beginning of this section.
One hundred parts of the earthy carbonates contain respectively, when quite pure, the following proportions of carbonic acid:
- Carbonate of magnesia, 52.38 carbonic acid. - Lime, 43.14 - Stroonian, 29.73 - Barytes, 22.00
Knowing these proportions, we may easily deduce from the weight of carbonic acid, obtained from any native carbonate subjected to analysis, the degree of its purity with very little trouble.
We might lay down here rules for analysing the other earthy salts, as sulphates, phosphates, tungstates, &c. which occur ready formed in the earth. But such details would swell this article to too great a length. We must refer to the observations which we have already made on the analysis of the salts. The young analyst, who wishes to become expert in these analyses, cannot do better than procure a copy of Klaproth's Beiträge, or Analytical Essays, in six octavo volumes. The two first of them have been translated into English; but the last four still remain in the original German. In these volumes he will find examples of the analysis of almost every earthy salt, and, by imitating the methods there laid down, he will soon become an expert analyst, and will be able even to improve upon Klaproth's processes, by applying to them various discoveries which have been made since the original publication of these experiments.
**CHAP. VI.—Of Ores.**
The term ore is applied to all those mineral bodies from which metals are extracted to answer the purposes of civilized society. The term has been extended, by mineralogists, to all those minerals which contain a notable proportion of metal, whether that metal be considered of such value as to be extracted from the ore or not. Thus, wolfram is considered as an ore, though the tungsten, which constitutes its principal constituent, has not hitherto been applied to any useful purpose, and, of course, though wolfram is never collected for the purpose of extracting from it the tungsten for the uses of civilized society.
The ores are so numerous, and so complex in their nature, that a general formula of analysis cannot be applied to them all. Many of them, indeed, cannot be looked upon in any other light than as mere mechanical mixtures; an accurate analysis of which cannot be of any utility for the purposes of science, because no two portions would yield exactly the same result; though it may frequently be of consequence to miners and metallurgists to know the constituents of even such mechanical mixtures, when they happen to occur abundantly in any peculiar place; because such knowledge will facilitate the invention of processes for extracting the useful metals out of the mixture.
To enter upon a minute detail of the various methods employed for the analysis of the ores, would be to extend this article to a disproportionate length. All we can do is, to make a few general observations on the mode of analysing the most important ores, which are likely to come in the way of the young chemist.
Metals occur in the earth in five different states,
1. As metals, either alone, or more commonly united with each other, constituting alloys. 2. United to sulphur, constituting sulphurets. 3. United to oxygen, constituting oxides. 4. United to chlorine, constituting chlorides. 5. In the state of oxides united to acids, and constituting salts.
As the mode of analysis differs materially according to these states, it will be proper to consider each of them separately.
I.—Alloys.
Gold, platinum, palladium, iridium, and tellurium, have hitherto been found only in the state of alloys.
Silver, copper, bismuth, and arsenic, occur very frequently in that state; and mercury, iron, cobalt, nickel, and antimony, occur in it occasionally.
Gold occurs either pure, or alloyed with silver or copper, or both. The mode of analysing such alloys has been given in a preceding chapter.
No accurate and simple mode of analysing the alloys of platinum, palladium, and iridium, has yet been found out. For the methods at present known, the reader is referred to Wollaston, Phil. Trans. 1804 and 1805; Tennant, Phil. Trans. 1804; and Vauquelin, Ann. de Chimie, 88.
Tellurium occurs alloyed with gold, iron, silver, and lead. If the alloy be treated with nitro-muriatic acid, all the metals will be dissolved, except the silver, which will remain in the state of a chloride, from which the quantity of it may be ascertained. Potash poured in excess into the liquid will precipitate the gold and the iron, but will keep the tellurium and the lead in solution. If the precipitate of oxides of gold and iron be moderately heated, the gold will be restored to the metallic state. Muriatic acid will then dissolve the iron, and leave the gold untouched. We may precipitate the iron by means of ammonia, and determine its quantity. The lead may be precipitated from the potash solution by means of sulphate of soda, and the tellurium by saturating the potash with muriatic acid.
Silver, when in the metallic state, occurs either pure, or alloyed with gold, antimony, lead, arsenic, bismuth, or iron. It may be dissolved in nitric acid,—the gold will remain in the metallic state. Water will precipitate the antimony and bismuth, sulphate of soda the lead, muriate of lime the arsenic, and ammonia the iron.
Native copper has scarcely been subjected to analysis. It may be dissolved in nitric acid. Sulphate of soda will precipitate the lead, if it contain any, and common salt the silver.
The other native metals have not hitherto been subjected to analysis. We may proceed according to the rules laid down in a preceding chapter, when treating of the metals.
II.—Sulphurets.
The following metals occur in the state of sulphuret: Mercury, silver, copper, iron, cobalt, molybdenum, tin, zinc, bismuth, lead, antimony, and arsenic.
In these sulphurets the proportion of sulphur and of metal is always such, that if the sulphur were converted into sulphuric acid, and the metal into an oxide, the acid and oxide would be capable of uniting together, and of forming a neutral sulphate. The method of analysis is, to treat the sulphuret with nitric acid, till the whole sulphur is acidified. In some cases, as mercury, silver, lead, we shall, by this process, form an insoluble sulphate. If this sulphate be washed, dried, and weighed, it will enable us to determine the proportion of the constituents of the sulphuret before the experiment. Thus, 100 parts of sulphuret of lead, when converted into sulphate of lead, will weigh $126\frac{3}{4}$; of this $\frac{4}{5}$ths, or 0.21, consist of oxygen added by the action of the nitric acid. The remaining $\frac{1}{5}$ths consist of sulphur and lead, in the proportion of 2 sulphur and 13 lead, or, in other words, sulphuret of lead is composed of
| Lead | 86$\frac{3}{4}$ | |------|--------------| | Sulphur | 13$\frac{1}{4}$ |
When sulphuret of bismuth or antimony have been converted into sulphates, if we pour water in sufficient quantity upon the mass, the acid is dissolved in the liquid, while the oxides of the metals remain in the state of a white powder. This powder being washed, dried, and heated to redness, will give the weight of the antimony or the bismuth respectively, by applying the rules laid down in a preceding chapter. The sulphuric acid is to be precipitated in the state of sulphate of barytes, and the weight of the acid is to be estimated by the rules previously given. Knowing the weight of sulphuric acid, it is easy to deduce that of the sulphur; for $\frac{3}{4}$ths of the weight of that acid are sulphur, and the remaining $\frac{1}{4}$ths oxygen.
III.—Oxides.
The following metals are found in the earth in the state of oxides: Silver, copper, iron, manganese, uranium, cerium, tantalum, cobalt, nickel, tin, titanium, zinc, bismuth, lead, antimony, and arsenic.
These oxides occur sometimes in a separate state, but more frequently mixed together, or united either mechanically or chemically to different earths, particularly alumina. To be able to analyse them, we must make ourselves well acquainted with the properties of the metals, as detailed in a preceding chapter, and with the various oxides which these metals are capable of forming; their solubility or insolubility in various acids; and the action of the alkalies and various saline solutions upon these liquids. Upon a knowledge of these properties depends the power of being able, successfully, to separate them from each other.
As an example of the mode of analysing a mixture of metallic oxides, we may take red silver ore, which has the aspect of a chemical compound of silver and antimony, both in the state of oxides, and containing likewise a quantity of sulphur. When digested in diluted nitric acid, about 58 per cent. of it is dissolved. When the residual parts are digested in muriatic acid, a part dissolves, and a portion remains undissolved. This last portion burns all away, with a blue flame, when heated, and consequently is sulphur. When the muriatic acid solution is diluted with water, a white powder falls, which is hydrate of antimony. The nitric acid solution, being mixed with common salt, the silver is precipitated in the state of chloride. These are the only constituents of the ore. We have given the analysis of red silver ore, because we are far from thinking that it has hitherto been analysed in a perfectly satisfactory manner. The present advanced state of the science would enable the chemist to go farther than Klaproth or Vauquelin were able to do at the time that their analyses were published, and to determine what portion of the respective metals are in the state of oxides, and what portions are combined with sulphur in the metallic state.
IV.—Chlorides.
The only metals hitherto found in the earth, in the state of chlorides, are mercury, silver, and lead; mercury, in the state long well known by the name of Calomel, which is a protochloride; and the silver and lead, in what was formerly called Horn-silver and Horn-lead. These chlorides are easily reduced to the metallic state, by means of hydrogen, or when treated with substances containing hydrogen. Knowing the weight of metal which they contain, it is easy to deduce the proportion of chlorine which makes up the remainder of their weight. The proportion of chlorine and metal which exists in each of these chlorides is as follows:
Chlorine: \[ 4.5 + 25 \text{ mercury} = 29.5 \text{ chloride of mercury} \] \[ 4.5 + 13.75 \text{ silver} = 18.25 \text{ chloride of silver} \] \[ 4.5 + 13 \text{ lead} = 17.5 \text{ chloride of lead} \]
V.—Salts.
The following metals are found in the earth in the state of salts: Silver, copper, iron, manganese, cerium, cobalt, titanium, zinc, and lead.
Silver occurs in the state of carbonate.
Copper in the state of carbonate, silicate, carbonatite, arseniate, muriate, phosphate, and sulphate.
Iron in the state of carbonate, phosphate, arseniate, chromate, silicate, tungstate, and sulphate.
Manganese in the state of phosphate and silicate.
Cerium in the state of silicate and fluate.
Cobalt in the state of arseniate and sulphate. Titanium in the state of silicate.
Zinc in the state of silicate, carbonate, and sulphate.
Lead in the state of carbonate, muriato-carbonate, phosphate, arsenio-phosphate, chromate, sulphate, molybdate, arseniate.
For the rules by which the analysis of these salts is to be conducted, we must refer to our observations on the mode of analysing salts, in a preceding chapter of this article.
**Chap. VII.—Of Vegetable Substances.**
The species of plants are so numerous, so different from each other, so beautiful, and so useful, that they are likely to present themselves to the chemist more frequently than any other class of bodies whatever. Unfortunately, our methods for analysing vegetable substances are much more imperfect than those which we employ in our experiments upon mineral bodies. Vegetable substances are so numerous, so easily altered or even destroyed, our processes for separating them from each other are so imperfect, that no two analyses of the same body could be expected to correspond exactly with each other. All that we can do, therefore, is to point out the processes which are usually followed, and the substances which these processes are capable of separating from each other. The greater number of vegetable substances either exude spontaneously from plants, or they are obtained by processes which sacrifice all the other constituents to one peculiar and important substance; which it is the object of the process to procure. Thus gums, resins, birdlime, caoutchouc, the gum-resins, and balsams, exude spontaneously. Sugar, indigo, starch, oils, both fixed and volatile; camphor, wax, &c., are obtained by simple and well known processes from the peculiar plants which yield them, while every other part of the plant is sacrificed to these favourite educts.
In general, it is not possible to make the weight of the substances extracted from any part of a plant to correspond with the weight of the part from which they were extracted. We have no mode of bringing vegetable bodies to the same accurate state of dryness. If we expose them to a red heat, they all undergo decomposition, and are converted into water, carbonic acid, carbureted hydrogen, oil, acetic acid, and other similar substances, quite different in their nature from the body which we exposed to heat. It is very seldom that we can venture to expose a vegetable substance to a temperature higher than that of boiling water. Some of them, as saccharine matter of malt, are not capable of bearing even that temperature without alteration. Nor can there be any reason to doubt that the leaves, the petals, and even the bark of many plants will be altered, if we expose them to such a temperature.
We are under the necessity, in consequence of this easy alterability of vegetable bodies, to subject them to analysis precisely in the state that we find them; though this precludes the possibility of making the weight of the educts tally with that of the substance subjected to analysis. These educts are dried at the temperature of boiling water when they are capable of bearing that temperature, and when this is not the case, we are obliged to weigh them as we get them.
There is another method of drying vegetable bodies, which has been put in practice of late years, and which is undoubtedly better than exposing them to heat. It consists in putting them in the exhausted receiver of an air-pump, while a flat glass dish, containing a quantity of concentrated sulphuric acid, covers the bottom of the receiver. By this contrivance, they are more thoroughly dried than they would be by the heat of boiling water, while they run no risk of being decomposed by heat. But even this method cannot be applied to the leaves, petals, and other soft and delicate parts of plants, without producing considerable alterations. The leaves and petals change their colour, and the more volatile parts of these organs make their escape, and are absorbed by the sulphuric acid. Upon the whole, therefore, it is better to take the parts of plants to be subjected to analysis just as we find them. We may form an estimate of the quantity of water which they contain, by heating another portion, or exposing it under an exhausted receiver along with sulphuric acid.
Let us suppose then a vegetable substance, as a wood, bark, or the leaves of some particular tree, to be subjected to analysis. In order to give a precise example of the methods followed, we shall take the root of the Glycyrrhiza glabra or liquorice, a plant which is cultivated in the south of Europe for the sake of the saccharine extract which is obtained from it, and which is well known by the name of liquorice sugar. This root has been frequently subjected to analysis by chemists, but the most instructive analysis of it which has yet appeared is by Robiquet, in the Annales de Chimie, Vol. LXXII., p. 143.
A quantity of fresh liquorice root, washed quite clean, and divided as much as possible, was put into cold distilled water, and allowed to remain for twelve hours, when the liquid was filtered. It had assumed a reddish brown colour, and of course had dissolved a portion of the root.
This liquid being left at rest for some time, allowed a white powder to precipitate, which possessed the properties of starch.
The liquid thus freed from starch had a sweet taste accompanied with a certain acidity. When examined by reagents, it exhibited the following properties:
1. It reddened paper stained with litmus, and of course contained an acid. 2. Infusion of nut-galls and alcohol occasioned a slight precipitate. Hence it probably contained vegetable albumen. 3. A solution of glue in water occasioned no change in it. Therefore, it contained no sensible portion of tannin. 4. The acids occasioned a copious coagulation in it. 5. Potash occasioned no other change but altering the shade of colours. Hence it contained no notable proportion of earthy salt. 6. Lime water occasioned a copious precipitate. Hence it might contain phosphoric acid.
7. Oxalate of ammonia occasioned a gritty white precipitate. This might indicate the presence of lime. But the presence of this earth is rendered unlikely, by the action of potash on the liquid.
8. Acetate of lead formed a very copious magma. This might be owing to the presence of malic or phosphoric acid, or several other vegetable substances with which the oxide of lead unites.
9. The salts of iron and the muriate of barytes likewise occasioned similar precipitates. This seems to indicate the presence of phosphoric acid in the liquid.
The coagulation produced by the acids indicating the presence of some peculiar principle in the liquid, the next object was to separate this principle, and examine its properties. For this purpose, a quantity of the liquid was heated in a glass vessel, and boiled for a few minutes, to separate the vegetable albumen, the presence of which had been indicated by the action of infusion of nut-galls. For vegetable albumen conglobates and precipitates in flocks, when the liquid holding it in solution is heated. The flocks which fell by this process being separated by the filter, the liquid was allowed to cool, and, when quite cold, a little distilled vinegar was mixed with it. Some flocks appeared at first, which speedily increased in quantity so much, that they exhibited the appearance of a transparent gelatinous magma, which, when separated by the filter, and washed with cold water, constituted the substance to which liquorice owes its sweet taste. This substance possesses peculiar characters, and approaches sarcocoll in its properties.
Its colour is yellow, its taste precisely similar to that of liquorice. Its bulk diminishes very much in drying. The dried mass is scarcely soluble in cold water, but it dissolves readily in hot water; and if the solution be concentrated, it assumes, on cooling, the state of a transparent solid jelly. Cold alcohol dissolves it readily, and assumes a dark yellow colour, a syrupy consistence, and a sweet taste. When dissolved in water, and mixed with yeast, it does not ferment; neither does it yield any oxalic or malic acid when treated with nitric acid; but is converted chiefly into a yellow opaque mass, having but little taste, and similar in appearance to a resin.
The liquid, which had been deprived of the saccharine matter by means of distilled vinegar, still retained its colour. Acetate of lead formed in it a copious precipitate, and left the liquid colourless. A portion of the leaden precipitate being exposed to heat before the blow-pipe blackened, and then fused into an opaque bead, which, on cooling, assumed the shape of an irregular polyhedron. This indicates that it consisted chiefly of phosphate of lead. The whole of the precipitate being diffused through water, the lead was thrown down by means of a current of sulphurated hydrogen gas. There remained in solution in the water phosphoric acid, malic acid, and the substance which gave the infusion of liquorice its yellowish brown colour.
The original liquid, which had been freed from its colour by means of acetate of lead, was treated with sulphurated hydrogen gas, to throw down any excess of lead which might still remain in it. It was then filtered, concentrated by evaporation, and set aside for some days. Fine transparent crystals were deposited, having the form of rectangular octahedrons. Their nature was not particularly examined, but they probably consisted of asparagin.
Such are the substances which may be extracted from liquorice root by means of water. The next step in the analysis of a vegetable substance, is to try the effect of alcohol upon it.
A quantity of liquorice root was macerated in successive portions of alcohol, till it ceased to communicate any colour to that liquid. When the alcoholic solution was evaporated, a thick brown viscoid oil separated from it. This oil had at first a sweetish taste, but it soon left a very acid impression, which was felt most strongly in the throat.
Liquorice root, having been thus treated by water and alcohol, was exposed to the action of very dilute nitric acid. But, after macerating in that liquid for a fortnight, nothing was dissolved but a little phosphate of lime.
The residual matter now consisted chiefly of woody fibre. To determine what saline bodies it contained, it was burnt. It left a bulky ash, which contained a good deal of carbonate of lime and phosphate of lime.
As a second example of vegetable analysis, we shall give that of rice, or the seeds of the Oryza sativa, as performed by Bracconnot (Ann. de Chim. et Phys. 4. 370).
One hundred parts of Carolina rice being dried, were found to lose 5 per cent. of their weight. Being put into water, of the temperature of 122°, they absorbed that liquid with avidity, and opened transversely in several parts. When so soft that they could easily be squeezed into a fine powder between the fingers, they were thoroughly pounded in a glass mortar, along with a sufficient quantity of the water in which they had been macerated. A milky liquid was thus obtained, which was thrown upon a filter. The greatest part of the substance of the rice remained upon the filter. This portion, being well washed and dried, weighed 93.67. The watery liquid which passed through the filter we shall call A.
The 93.67 parts, which remained upon the filter, being diffused through water, passed entirely through a silk scirse. This milky liquor contained two distinct substances: the first, which was very white, and by far the most abundant, remained long in suspension; but the second, which was much less white, was of a greater specific gravity, and, of course, fell more speedily to the bottom. Advantage was taken of this difference, to separate these two substances from each other. The whitest of the two possessed all the properties of starch. It had a fine white colour, was light, and, when pressed, a peculiar noise could be perceived. When triturated in a mortar with a little iodine, it assumed a deep blue colour. It formed a kind of jelly when boiled in water. When one part of it was boiled in 4000 parts of water, the liquid, which was limpid, was precipitated by lime water, barytes water, and the infusion of nut-galls. The other substance was the parenchyma of the rice, but not quite free from starch. The liquid A, having the property of reddening vegetable blues, was put into a retort, and a portion of it distilled over. This portion was mixed with a little barytes water, and then evaporated to dryness. A small portion of residue remained, which, being placed in contact with diluted sulphuric acid, gave out the odour of acetic acid. Hence it is evident, that the liquid A owed the property of reddening vegetable blues to the presence of a little acetic acid.
During the distillation of the liquid A it became opaque, and at last let fall a small quantity of white matter, which had not the aspect of albumen. The whole of the liquid, including the precipitate, was poured into a small porcelain capsule, and evaporated to dryness. The residue weighed 1.28; it had a pale yellow colour, and attracted humidity from the atmosphere. It was mixed with a little water to give it the consistence of a syrup, then alcohol was poured into it. A copious precipitate fell, which, being collected and dried, weighed 0.99. This matter dissolved in cold water, with the exception of 0.13 of white floes which remained undissolved, and possessed the properties of albumen. The dissolved portion approached more nearly to the properties of torrefied starch than to those of gum.
The alcohol held in solution 0.29 of a syrupy matter, having the taste of sugar, the smell of honey, incapable of crystallizing and attracting humidity from the atmosphere.
The gummy matter contained mixed with it phosphate of lime, phosphate of potash, and a little lime combined with a vegetable acid.
The parenchyma was completely freed from starch, by boiling it for half an hour in water acidulated with sulphuric acid. The weight of the starch was found to be 85.07; that of the parenchyma 4.8.
When macerated rice is digested in alcohol for 24 hours, if the alcohol be separated by the filter and evaporated, it leaves behind a small quantity of a fat oil, almost colourless, having a rancid taste and smell, the consistence of thick olive oil, and easily soluble in alcohol and alkalies.
These examples are sufficient to give the reader some notion of the method of proceeding, when vegetable bodies are to be analysed. The art has not yet reached such perfection as to enable us to lay down general formulas, which will serve for every case. Almost always we begin by treating the vegetable substance with water. When every thing soluble in water has been removed, we treat it with alcohol. After the action of these two solvents has been exhausted, we sometimes employ sulphuric ether, sometimes acetic acid, sometimes dilute nitric acid, and sometimes potash ley, according to the nature of those substances which we suppose to exist in it. The young chemist who intends to devote himself to vegetable analysis, ought, in the first place, to make himself familiarly acquainted with the properties of all the vegetable principles, with their solvents and their precipitants, by a set of experiments, which must be repeated with all the requisite accuracy, till the facts are engraved upon his memory. He may then repeat some of the most elaborate vegetable analyses that have hitherto been performed.
For a knowledge of these analyses we refer him to the 4th volume of Dr Thomson's System of Chemistry, p. 207 (5th edition), where he will find a pretty complete collection of them, with exact references to the original books, where these analyses were published.
But the vegetable principles themselves are all compounds of various proportions of oxygen, carbon, hydrogen, and azote, and, in order to form correct notions respecting their nature and constitution, it is requisite to be acquainted with the exact proportions in which their constituents exist in each of them. It is only since the introduction of the Atomic Theory into chemistry, that such investigations can be attended with much advantage. But an accurate knowledge of the composition of vegetable bodies would undoubtedly throw considerable light on the atomic theory, which is still in a very imperfect state. It would be requisite, however, to apply the atomic theory to vegetable analysis, with much greater caution than has been hitherto done. For when we find by analysis that any vegetable principle yields a certain number of atoms of hydrogen, carbon, and oxygen, we have no right to infer that these atoms were immediately united with each other. They might have been previously united into a certain number of binary or even ternary compounds, and these compounds, united together, may constitute the vegetable substance which we have subjected to analysis. Our object at present is merely to point out the method of determining the proportion of oxygen, carbon, and hydrogen, which exist in any vegetable principle, as sugar, gum, starch, &c.
The first successful attempt at such analysis was made by Gay-Lussac and Thenard in the second volume of their Memoires Physico-Chimique. They made up the vegetable substance to be analysed into balls, with a certain quantity of chlorate of potash; taking care to mix them as intimately as possible. These balls were burnt in a glass vessel by the application of the heat of a lamp. The carbon was converted into carbonic acid gas, the hydrogen into water. The quantity of carbonic acid and water yielded by a determinate weight of any vegetable principle being known, it was easy to determine the weight of carbon and hydrogen which it contained, and the deficit of weight was considered as oxygen. This method does not answer for those vegetable substances that contain azote as a constituent. For in such cases ammonia or nitric acid might be formed, which would prevent the possibility of deducing any consequence from the experiment.
The method of Gay-Lussac and Thenard was considerably modified by Berzelius, in his paper on the analysis of vegetable substances, published in the 4th volume of Thomson's Annals of Philosophy. He employed a mixture of chlorate of potash and the vegetable body, as the French chemists had done; but the combustion took place in a glass tube hermetically sealed to another glass tube, containing muriate of lime, and from this passed another tube into the mercurial trough. The water formed was condensed in the muriate of lime, and its weight... ascertained. The carbonic acid gas passed into a glass jar standing over mercury, and its bulk and purity were easily determined. This method had an obvious advantage over that of the French chemists, who had no means of collecting and estimating the water formed during the process. But it is obviously impossible by means of Berzelius's method, any more than by that of the French chemists, to determine the composition of those bodies which contain azote.
More lately, Gay-Lussac has suggested a new method of analysing vegetable substances; namely, heating them in a tube with peroxide of copper. This method we have frequently practised. It seems easier than either of the two preceding ones, and has the advantage of enabling us to determine the composition of bodies which contain azote. For if the action of the heat be properly regulated, the azote comes over in the state of gas, mixed with the carbonic acid gas, formed by the process, and the quantity of each of these gases is easily determined by analysis. The water is collected by means of a tube filled with dry muriate of lime. We shall, therefore, describe here the method of analysing vegetable bodies by means of peroxide of copper.
Procure a copper tube, bored from a solid copper rod about twelve inches long, and with a bore of about one-third of an inch in diameter. To the mouth of this tube a brass tube must be ground air-tight. This brass tube may be about four inches long, and bent as in figure 22. The end a of the tube being the one which is ground to the mouth of the copper tube. To the end b of the brass tube, a glass about six inches long is fitted, so as to be nearly air-tight. This glass tube is bent at the end, so that it can be introduced below the mouth of a glass jar, standing inverted upon the mercurial trough and full of mercury. The glass tube is filled with dry muriate of lime, in the state of powder. The upper and lower ends are filled with amianthus. Its weight must be carefully ascertained and written down. For greater security, let the glass tube be luted to the brass tube. Fill the brass tube with amianthus. Weigh out three grains of the vegetable substance to be analysed. It is best to take it in its natural state, without freeing it from the water which it may contain. This water must be ascertained by other experiments, and its quantity allowed for in the analysis. Mix these three grains with 120 grains of peroxide of copper, previously reduced to the state of a fine powder. The mixture must be as intimate as possible, that is to say, the vegetable substance must be equally diffused through the whole of the oxide of copper. Put the mixture into the copper tube. It will fill about five inches of it. Fill up the tube now completely to the mouth with peroxide of copper, and put a little amianthus over it, to prevent any of the oxide from falling out. Then fix the brass tube in the copper one. Put the copper tube upon a small iron chaffer or cradle, so that one-half of the tube is within the chaffer and the other half on the outside, and the whole apparatus must be so placed, that the extremity of the glass tube is below an inverted glass jar, standing over mercury, in the mercurial trough. That portion of the copper tube which is on the outside of the chaffer, is now to be covered with a coat of moist clay, about an inch in thickness. This covering will prevent the heat from passing nearly so rapidly along the copper tube; the consequence of which will be, that the brass tube will remain comparatively cool during the whole experiment. A few pieces of burning charcoal are now placed round the portion of the copper tube which is within the chaffer. The fire is made to commence at the end of the tube nearest the mercurial trough, and it proceeds gradually backwards to the bottom of the tube. Care should be taken to keep the fire low, and to let the combustion proceed slowly. The copper tube need hardly be heated red-hot; though, to be sure that the combustion has been complete, we are always in the habit, just before we terminate the process, to make the whole of the copper tube within the chaffer distinctly red-hot. If the vegetable substance analysed happens to contain azote, there seems to be a temperature somewhere about a red heat, at which nitric acid is formed. As such a product would destroy the accuracy of the experiment, it is material to keep the heat so low as to prevent the risk of any such formation.
As soon as the temperature rises sufficiently high, the vegetable substance is completely decomposed. The carbon which was contained in it combines with oxygen, contained in the oxide of copper, and is converted into carbonic acid. The hydrogen combines with oxygen, and is converted into water, while the azote makes its escape in the gaseous state. The carbonic acid and azotic gases will be collected in the graduated glass jar over the mercury. To ascertain the bulk of each, let up into the jar a quantity of potash ley, and let it stand in contact of the gases for twenty-four hours. The whole of the carbonic acid will be absorbed. The diminution of bulk will give the quantity of carbonic acid, while the residual bulk will give the quantity of azotic gas; making allowance for the alteration in the bulk occasioned by the column of mercury in the jar, and by any change of temperature that may have taken place during the continuance of the experiment. Indeed, it is always necessary to reduce the gases to the bulk which they would occupy at the temperature of 60°, and when the height of the barometer is 30 inches;—because the specific gravity of these gases, which enters as an indispensable element in our calculations, was estimated at that temperature. Knowing the bulk of these gases, it is easy to deduce their weight, and hence to know the quantity of carbon and azote which the three grains of the vegetable substance analysed contained. The increased weight of the muriate of lime will give us the water formed by the process, or separated from the vegetable substance by the heat. Subtracting from this weight the known proportion of water contained in the vegetable substance, determined by other experiments, the remainder will be the water formed by the union of the hydrogen in the vegetable substance with the oxygen of the oxide. One-ninth of the weight of this water is the hydrogen contained in the three grains of vegetable substance analysed. Thus we determine the weight of the azote, carbon, and hydrogen, which our vegetable substance contains. Add all these weights together. If they amount to three grains, we may conclude that our substance contains no oxygen; but if the weight (as will almost always happen) be less than three grains, we must suppose that the substance contained a quantity of oxygen, the weight of which, when added to that of the other constituents, will make up the weight of three grains (the water which exists as a constituent of the body being supposed subtracted).
It may be worth while to give an example or two of this mode of analysis, that the reader may be able to form a correct idea of it.
0.333 part of acetic acid united to protoxide of lead, furnished, when burnt by means of chlorate of potash, 0.18 water, and 0.574 of carbonic acid. From this result it is inferred by Berzelius, that 100 of acetic acid are composed of:
| Hydrogen | 6.35 | |----------|------| | Carbon | 46.83| | Oxygen | 46.82|
100.00
Four grains of uric acid being mixed with peroxide of copper, and analysed in the way above described by Dr Prout, yielded
Water, - 1.05 grain. Carbonic acid gas, 11 cubic inches. Azotic gas, - 5.5 ditto.
Hence he infers that it is composed of
| Hydrogen | 0.11 or 2.857 | |----------|---------------| | Carbon | 1.37 | | Azote | 1.61 | | Oxygen | 0.91 |
4.00 100.000
Five grains of crystallized triple prussiate of potash yielded to Dr Thomson, when treated in the same way,
Water, 1.55 grain. Carbonic acid, 5.205 cubic inches. Azotic gas, 2.420 ditto.
Now, the gaseous part of the acid in five grains of this salt (excluding the iron which it contains) weighs 1.54 grains. Hence it is concluded that this gaseous part is composed of
| Carbon | 0.6579 or 42.51 | |--------|-----------------| | Azote | 0.7175 | | Hydrogen | 0.1722 |
But the mere knowledge of the respective weights of the elements of which these bodies are composed, throws but little light on their constitution. We must know the number of atoms of each of the elements which exist in them. Now, the weight of an atom of each of these elementary bodies has been ascertained to be as follows:
Oxygen, 1.00 Azote, 1.75 Carbon, 0.75 Hydrogen, 0.125
The comparison of the weights with the weights of the elements found in one hundred parts of the vegetable body analysed, will give the relative proportion of atoms of each. Thus, in the case of acetic acid, the numbers
6.35 46.83 46.82
bear the same ratio to each other, as the numbers
0.125 × 3 0.75 × 4 1.00 × 3
That is to say, that acetic acid may be conceived a compound of 3 atoms hydrogen + 4 atoms carbon, + 3 atoms oxygen.
But before we can know whether these are the true number of atoms in acetic acid, we must know what the relative weight of an atom of acetic acid is. This is known when we ascertain the weight of acetic acid which is capable of saturating an atom of each of the bases. Now, according to the analysis of Berzelius, acetate of lime and acetate of lead are composed as follows:
Acetate of lime. Acetic acid, 100 or 6.615 Lime, 54.8 3.625
Acetate of lead. Acetic acid, 100 or 6.432 Oxide of lead, 217.662 14.
Now, as these salts are neutral, we have reason to conclude that they are composed each of an atom of acetic acid, united to an atom of base. An atom of lime weighs 3.625, and an atom of oxide of lead 14. We see from the first salt that an atom of acetic acid weighs 6.615, and from the second that it weighs 6.432. The second of these determinations is more likely to be correct than the first, because it is easier to free acetate of lead of its water than acetate of lime. Hence we may conclude, that the weight of an atom of acetic acid does not exceed 6.432. Let us now see what the weight of 3 atoms hydrogen, 4 atoms carbon, and 3 atoms oxygen will be.
3 atoms hydrogen = 0.125 × 3 = 0.375 4 atoms carbon = 0.75 × 4 = 3.000 3 atoms oxygen = 1. × 3 = 3.000
6.375
We perceive that they amount to 6.375. Now, as 6.375 is very near 6.432, but rather under it, we have a right to conclude that acetic acid consists of
3 atoms hydrogen. 4 atoms carbon. 3 atoms oxygen.
The small difference is either owing to errors in the Decomposition of acetate of lead, or to the impossibility of freeing that salt entirely from water.
**Chap. VIII.—Of Animal Substances.**
The analysis of animal substances is conducted nearly on the same principles as that of vegetable substances. But greater progress has been made in it. This is partly owing to the comparatively small number of animal principles hitherto discovered, at least in the bodies of the larger animals, to which chemical analysis has been hitherto chiefly confined. But it is mainly owing to the great importance which medical men have attached to an accurate knowledge of the constituent parts of the human body. This has occasioned a more laborious examination, and has led to pretty precise methods of detecting the presence of the different animal bodies, and of determining the weight of each.
The animal substances of which the soft parts of the larger animals are composed, and which exist chiefly in the liquids of the body, may be reduced nearly to the four following:
- Gelatin. - Albumen. - Fibrin. - Colouring matter of blood.
It will be requisite to give the characteristic properties of each of these bodies, and to show how their presence may be detected in such animal liquids as happen to contain them.
1. As far as we know at present, gelatin never exists as a constituent of any of the animal fluids. It is extracted from the skin and from the membranes by long boiling. They gradually dissolve in the water, and seem to be converted into gelatin.
Gelatin, when in a solid state, is the substance called in common language glue. When pure it is semitransparent and white. When put into cold water it swells very much, and assumes a gelatinous consistence, but does not dissolve. In hot water it dissolves; but when the liquid cools it becomes a solid jelly, provided the gelatin dissolved in it amounts to \( \frac{1}{105} \)th of the weight of the water. One part of gelatin dissolved in 150 parts of hot water forms a liquid which does not gelatinize on cooling. This liquid is not precipitated by corrosive sublimate; but it is abundantly precipitated by infusion of nut-galls.
2. Albumen is the substance which gives the characteristic properties to the white of an egg and to the serum of blood. When either of these liquids is exposed to the temperature of 165°, it coagulates into a white or pearl coloured stiff mass. It is this coagulation by heat which characterizes albumen. If an acid, or alcohol, or sugar, or a salt, be dissolved in a liquid containing albumen, coagulation takes place. A liquid containing albumen is precipitated copiously by corrosive sublimate and acetate of lead. The infusion of nut-galls throws it down likewise, but is by no means so delicate a test of albumen as it is of gelatin. For example, it does not detect one part of albumen dissolved in 1000 parts of water.
3. Fibrin is so called because the fibres of the muscles are chiefly composed of it. Blood also contains it as an essential constituent. When a quantity of blood newly drawn from an animal is allowed to remain for some time at rest, a thick red clot gradually forms in it and subsides. If we take out this clot and wash it gently in water till that liquid ceases to become coloured by it, the portion which remains undissolved is fibrin.
Fibrin thus obtained is a white substance, at first soft and elastic; but it becomes hard when dried. It is insoluble in cold water; but when that liquid is long boiled upon it, it becomes milky. Infusion of nut-galls throws down floccs from this solution, which do not adhere together like the tannate of glue. The fixed alkaline leys dissolve fibrin when assisted by heat; so does acetic acid. Muriatic acid combines with it in two proportions. That which contains the least acid dissolves in water; but the other, containing a greater proportion, is insoluble in that liquid. Alcohol and ether gradually convert fibrin into a kind of fatty matter.
4. The water with which the clot of the blood has been washed (supposing it previously, by means of blotting paper, freed from the serum) contains in solution the colouring matter of the blood; and this matter may be obtained in a separate state by coagulating that liquid by exposure to heat. When the coagulum is dried at the temperature of 158° it becomes black, hard, difficulty pulverizable, and breaks with a vitreous fracture. The action of reagents on it in this state is nearly the same as upon fibrin.
By these properties may these different substances be recognised in the animal bodies that contain them. The albumen, fibrin, and colouring matter undergo certain changes in the fluids of secretion, to fit them for the purposes for which they are intended. In the bile they are converted into picromel, in the milk to curd, sugar of milk and cream. In the secretion of the nose to mucus. Urine contains a peculiar substance, called urea, seemingly formed in the kidneys, and immediately thrown out of the body.
We had intended to have laid down formulas for the analysis of the different fluids of animal bodies. But this article has already extended to such a length, that we conceive it necessary to bring it to a termination. We would recommend those persons who wish to acquire skill in this kind of analysis, to study a work on the subject published by Berzelius in 1806 and 1808, in two octavo volumes, and entitled *Föreläsningar i Djurkemien*. The rules of animal analysis are laid down in that book with all the requisite minuteness, and excellent examples are given by the accurate analysis of most of the human secretions and excretions. It is a serious injury to the progress of animal chemistry in this country, that this work has not been translated into the English language.
None of the readers of this article can be more sensible of its imperfections than the author of it himself. But it was impossible to remedy these imperfections without entering into details quite incompatible with the limits of a supplementary article. Instead of a minute description of the method to be followed in the analysis of every kind of substance, which would have been requisite in a complete treatise on this subject, we have been obliged to satisfy ourselves with general observations on the mode of investigating some of the more prominent groups.
We are not without hopes that these observations, limited and imperfect as they are, will be of some use to the young chemist; especially those that refer to the gases and salts, which we have treated somewhat more at length than any of the other departments.