FOOTNOTES:

Mixed chlorides0.0266 gramDeduction0.0110 "———Mixed alkalies0.0156 "

Assuming this to have been got from 1 gram of a rock, it would amount to 1.56 per cent. of "potash and soda."

The relative proportions of the potash and soda can be ascertained from the same determination. Sodium and potassium chlorides have the following composition:—

Sodium39.38Potassium52.46Chlorine60.62Chlorine47.54——————100.00100.00

The percentage of chlorine in the mixed chlorides is calculated. It necessarily falls somewhere between 47.5 and 60.6 per cent., and approaches the one or the other of these numbers as the proportion of the sodium or potassium preponderates. Each per cent. of chlorine in excess of 47.5 represents 7.63 per cent. of sodium chloride in the mixed chlorides. The percentage of potash and soda in the substance can be calculated in the usual way. Sodium chloride multiplied by 0.5302 gives its equivalent of soda (Na2O), and potassium chloride multiplied by 0.6317 gives its equivalent of potash (K2O).

The weight of sodium chloride in the mixed chlorides is also calculated thus:—Take the same example for illustration. Multiply the chlorine found by 2.103. This gives—

(0.0142×2.103) = 0.02987.

From the product deduct the weight of the mixed chlorides found—

Product0.02987Mixed chlorides0.02660———Difference0.00327

The difference multiplied by 3.6288 gives the weight of sodium chloride in the mixture. In this case it equals 0.0118 gram. The potassium chloride is indicated by the difference between this and the weight of the mixed chlorides. It equals 0.0148 gram. We have now got—

Sodium chloride0.0118 gramPotassium chloride0.0148 "

from 1 gram of the rock taken. Multiplying these by their factors we have (Soda = 0.0118×0.5302; Potash 0.0148×0.6317)—

Soda= 0.625 per cent.Potash= 0.935 "

Concentration of the Alkalies.—With the exception of magnesia, all the other bases are separated from the alkalies in the ordinary course of work without the addition of any re-agent which cannot be removed by simple evaporation and ignition. Consequently, with substances soluble in acids, successive treatment of the solution with sulphuretted hydrogen, ammonia,ammonic sulphide, and ammonic carbonate, filtering, where necessary, will yield a filtrate containing the whole of the alkalies with ammonic salts and, perhaps, magnesia.

The filtrate is evaporated in a small porcelain dish, with the addition of nitric acid towards the finish. It is carried to dryness and ignited. The residue is taken up with a little water, treated with a few crystals of oxalic acid, and again evaporated and ignited. The alkaline salts are extracted with water, and filtered from the magnesia into a weighed platinum dish. The solution is then evaporated with an excess of hydrochloric acid, ignited at a low red heat, and weighed. The residue consists of the mixed alkaline chlorides.

For substances (such as most silicates and similar bodies) not completely decomposed by acids, Lawrence Smith's method is generally used. This is as follows:—Take from 0.5 to 1 gram of the finely powdered mineral, and mix, by rubbing in the mortar, with an equal weight of ammonium chloride. Then mix with eight times as much pure calcium carbonate, using a part of it to rinse out the mortar. Transfer to a platinum crucible, and heat gently over a Bunsen burner until the ammonic chloride is decomposed (five or ten minutes). Raise the heat to redness, and continue at this temperature for about three quarters of an hour. The crucible must be kept covered. Cool, and turn out the mass into a 4-inch evaporating dish; wash the crucible and cover with distilled water, and add the washings to the dish; dilute to 60 or 80 c.c., and heat to boiling. Filter and wash. Add to the filtrate about 1.5 gram of ammonium carbonate; evaporate to about 40 c.c., and add a little more ammonic carbonate and some ammonia. Filter into a weighed platinum dish, and evaporate to dryness. Heat gently, to drive off the ammonic chloride, and ignite to a little below redness. Cool and weigh. The residue consists of the mixed alkaline chlorides.

Separation of the Alkali-Metals from each other.—Sodium and lithium are separated from the other alkali-metals by taking advantage of the solubility of their chlorides in the presence of platinic chloride; and from one another by the formation of an almost insoluble lithic phosphate on boiling with a solution of sodium phosphate in a slightly alkaline solution. Cæsium, rubidium, and potassium yield precipitates with platinic chloride, which are somewhat soluble, and must be precipitated from concentrated solutions. Cæsium and rubidium are separated from potassium by fractional precipitation with platinum chloride. Their platino-chlorides, being less soluble than that of potassium, are precipitated first. One hundred parts of boiling water dissolve 5.18 of the potassium platino-chloride, 0.634 of the rubidiumsalt, and 0.377 of the corresponding cæsium compound. The separation of lithium, cæsium, and rubidium is seldom called for, owing to their rarity. The details of the separation of potassium from sodium are described underPotassium. Ammonia compounds are sharply marked off from the rest by their volatility, and it is always assumed that they have been removed by ignition; if left in the solution, they would count as potassium compounds. They will be considered underAmmonia.

Sodium is the commonest of the alkali metals. It is found in nature chiefly combined with chlorine as "common salt" (NaCl). This mineral is the source from which the various compounds of sodium in use are prepared. Sodium occurs abundantly as nitrate (NaNO3) in Chili saltpetre, and as silicate in various minerals, such as albite (or soda-felspar).

It occurs as fluoride in cryolite (Na3AlF6), and as carbonate in natron, &c. Sulphates are also found. Sodium is very widely diffused, few substances being free from it.

The detection of sodium is easy and certain, owing to the strong yellow colour its salts impart to the flame; this, when viewed by the spectroscope, shows a single yellow line.[93]The extreme delicacy of this test limits its value, because of the wide diffusion of sodium salts. It is more satisfactory to separate the chloride, which may be recognised by its taste, flame coloration, fusibility, and negative action with reagents. The chloride dissolved in a few drops of water gives with potassium metantimoniate, a white precipitate of the corresponding sodium salt.

Sodium salts are dissolved out from most compounds on treatment with water or dilute acids. Insoluble silicates are decomposed and the alkali rendered soluble by Lawrence Smith's method, which has just been described. The separation of the sodium from the mixed chlorides is effected in the following way:—The chlorides are dissolved in a little water and the potassium separated as platino-chloride. The soluble sodium platino-chloride, with the excess of platinum, is boiled, mixed with sulphuric acid, evaporated to dryness, and ignited. On extracting with water, filtering, evaporating, and igniting, sodium sulphate is left, and is weighed as such.

It is more usual, and quite as satisfactory, to calculate the weight of the sodium chloride by difference from that of the mixed chlorides, by subtracting that of the potassium chloride,which is separately determined. For example, 1 gram of a rock gave—Mixed chlorides, 0.0266 gram, and 0.0486 gram of potassic platino-chloride. This last is equivalent to 0.0149 gram of potassium chloride.

Mixed chlorides found0.0266Deduct potassium chloride0.0149———Leaves sodium chloride0.0117

The weight of sodium chloride found, multiplied by 0.5302, gives the weight of the soda (Na2O).

The solution, which must contain no other metal than sodium, is evaporated in a weighed platinum crucible or dish. Towards the finish an excess, not too great, of sulphuric acid is added, and the evaporation is continued under a loosely fitting cover. The residue is ignited over the blowpipe, a fragment of ammonic carbonate being added towards the end, when fumes of sulphuric acid cease to be evolved. This ensures the removal of the excess of acid. The crucible is cooled in the desiccator, and weighed. The substance is sulphate of soda (Na2SO4), and contains 43.66 per cent. of soda (Na2O), or 32.38 per cent. of sodium (Na).

There are various methods used for the different compounds of sodium. There is no one method of general application. Thus with "common salt" the chlorine is determined volumetrically; and the sodium, after deducting for the other impurities, is estimated by difference.

With sodic carbonate and caustic soda, a given weight of the sample is titrated with standard acid, and the equivalent of soda estimated from the alkalinity of the solution.

With sodium sulphate, a modification of the same method is used. To a solution of 3.55 grams of the salt contained in a half-litre flask, 250 c.c. of a solution of baryta water is added. The volume is made up to 500 c.c. with water. The solution is mixed and filtered. Half of the filtrate is measured off, treated with a current of carbonic acid, and then boiled. It is transferred to a half-litre flask, diluted to the mark, shaken up, and filtered. 250 c.c. of the filtrate, representing a quarter of the sample taken, is then titrated with standard acid. The standard acid is made by diluting 250 c.c. of the normal acid to 1 litre. The c.c. ofacid used multiplied by 2 gives the percentage. A correction must be made to counteract the effect of impurities in the baryta as well as errors inherent in the process. This is small, and its amount is determined by an experiment with 3.55 grams of pure sodium sulphate.

Moisture.—Powder and weigh up 10 grams of the sample into a platinum dish. Dry in a water oven for an hour, and afterwards heat to bare redness over a Bunsen burner. Cool, and weigh. The loss gives the water.

Chlorine.—Weigh up two separate lots of 1 gram each; dissolve in 100 c.c. of water, and determine the chlorine by titrating with the standard silver nitrate solution, using chromate of potash as indicator. SeeChlorine.

Insoluble Matter.—Dissolve 10 grams of the salt in water with the help of a little hydrochloric acid. Filter off the sediment, wash, ignite, and weigh. This residue is chiefly sand. Dilute the nitrate to 500 c.c.

Lime.—Take 250 c.c. of the filtrate, render ammoniacal and add ammonium oxalate; wash, dry, and ignite the precipitate. Weigh as lime (CaO).

Magnesia.—To the filtrate from the lime add phosphate of soda. Allow to stand overnight, filter, wash with dilute ammonia, dry, ignite, and weigh as pyrophosphate.

Sulphuric Oxide.—To the remaining 250 c.c. of the filtrate from the "insoluble," add an excess of barium chloride. Collect, wash, dry, ignite, and weigh the barium sulphate.

Sodium.—It is estimated by difference.

The following may be taken as an example:—

Moisture0.35Insoluble matter0.40Lime0.40Magnesia0.05Sulphuric oxide0.60Chlorine59.60Sodium38.60———100.00

Potassium occurs in nature as chloride, in the mineral sylvine (KCl), and more abundantly combined with magnesium chloride, in earnallite(KCl.MgCl2.6H2O). It occurs as nitrate in nitre (KNO3), and as silicate in many minerals, such as orthoclase (or potash-felspar) and muscovite (or potash-mica).

Potassium compounds are detected by the characteristic violet colour they impart to the flame. The presence of sodium salts masks this tint, but the interference can be neutralised by viewing the flame through a piece of blue glass. Viewed through the spectroscope, it shows a characteristic line in the red and another in the violet. These, however, are not so easy to recognise or obtain as the sodium one. Concentrated solutions of potassium salts give a yellow crystalline precipitate with platinum chloride, and a white crystalline one with the acid tartrate of soda. For these tests the solution is best neutral. These tests are only applicable in the absence of compounds other than those of potassium and sodium.

This process serves for its separation from sodium. Take 1 gram of the sample and dissolve it in an evaporating dish with 50 c.c. of water. Acidify with hydrochloric acid in quantity sufficient (if the metals are present as chlorides) to make it acid, or, if other acids are present, in at least such quantity as will provide the equivalent of chlorine. Add 3 grams of platinum, in solution as platinum chloride, and evaporate on a water-bath to a stiff paste, but not to dryness. Moisten with a few drops of platinic chloride solution without breaking up the paste by stirring. Cover with 20 c.c. of strong alcohol, and wash the crystals as much as possible by rotating the dish. Allow to settle for a few moments, and decant through a filter. Wash in the same way two or three times until the colour of the filtrate shows that the excess of the platinum chloride used is removed. Wash the precipitate on to the filter with a jet of alcohol from the wash-bottle; clean the filter-paper, using as little alcohol as possible. Dry in the water-oven for an hour. Brush the precipitate into a weighed dish, and weigh it. It is potassium platino-chloride (K2PtCl6), and contains 16.03 per cent. of potassium, or 30.56 per cent. of potassium chloride (KCl), which is equivalent to 19.3 per cent. of potash (K2O).

If the filter-paper is not free from precipitate, burn it and weigh separately. The excess of weight over that of the ash will be due to platinum and potassic chloride (Pt and 2KCl). This multiplied by 1.413 will give the weight of the potassic platino-chloride from which it was formed. It must be added to the weight of the main precipitate.

The mixed alkaline chlorides obtained in the usual course ofanalysis are treated in this manner; the quantity of platinum added must be about three times as much as the mixed chlorides weigh.

These are the same as with soda.

Examination of Commercial Carbonate of Potash.—The impurities to be determined are moisture, silica, and insoluble matter, chlorine, sulphuric oxide, and oxide of iron. These determinations are made in the ways described under the examination of common salt.

Thepotassiumis determined by converting it into chloride and precipitating with platinum chloride, &c., as just described.

Available Alkali.—Weigh up 23.5 grams of the sample, dissolve in water, and make up to 500 c.c. Take 50 c.c., tint with methyl orange, and titrate with the normal solution of acid. The c.c. of acid used multiplied by 2 gives the percentage of available alkali calculated as potash (K2O).

Soda.—This is calculated indirectly in the following way:—Deduct from the potassium found the quantity required for combination with the chlorine and sulphuric oxide present, and calculate the remainder to potash (K2O). The apparent surplus excess of available alkali is the measure of the soda present.

Carbon Dioxide.—The c.c. of acid used in the available alkali determination, multiplied by 2.2 and divided by 2.35, gives the percentage of carbon dioxide.

Lithia, the oxide of lithium (Li2O), occurs in quantities of 3 or 4 per cent. in various silicates, such as lepidolite (or lithia-mica), spodumene, and petalite. It also occurs as phosphate in triphyline. It is a constituent of the water of certain mineral springs. A spring at Wheal Clifford contained as much as 0.372 gram of lithium chloride per litre. In small quantities, lithia is very widely diffused.

TheDetectionof lithia is rendered easy by the spectroscope; its spectrum shows a red line lying about midway between the yellow sodium line and the red one of potassium. It also shows a faint yellow line. The colour of the flame (a crimson) is characteristic.

The reactions of the lithium compounds lie between those of the alkalies and of the alkaline earths. Solutions are not precipitated by tartaric acid nor by platinic chloride. The oxide isslowly soluble in water. The carbonate is not freely soluble. Lithia is completely precipitated by sodic phosphate, especially in hot alkaline solutions.

In its determination the mixed alkaline chlorides obtained in the separation of the alkalies are dissolved in water, a solution of soda is added in slight excess, and the lithia precipitated withsodicphosphate. Before filtering, it is evaporated to dryness and extracted with hot water rendered slightly ammoniacal. The residue is transferred to a filter, dried, ignited, and weighed. The precipitate is lithium phosphate (3Li2O, P2O5), and contains 38.8 per cent. of lithia. The separation of lithia from magnesia is not given by the usual authorities. Wohler recommends evaporating the solution to dryness with carbonate of soda. On extracting the residue with water, the lithia dissolves out and is determined in the filtrate. One hundred parts of water dissolve, at the ordinary temperature, 0.769 parts of lithium carbonate (Li2CO3); the basic magnesia compound is almost insoluble in the absence of carbon dioxide and ammonium salts.

The oxide of caesium, caesia (Cs2O), is found associated with lithia in lepidolite, &c., and, together with rubidium, in many mineral waters. The mineral pollux is essentially a silicate of alumina and caesia; it contains 34.0 per cent. of the latter oxide.

Caesium is best detected by the spectroscope, its spectrum being characterised by two lines in the blue and one in the red; the latter is about midway between the lithium and sodium lines.

If not detected by the spectroscope, or specially looked for, caesia would, in the ordinary course of work, be separated with the potash and weighed as potassium platino-chloride.

Caesia is separated from all the other alkalies by adding to the acid solution of the mixed chlorides a strongly acid cold solution of antimonious chloride. The acid used must be hydrochloric. The caesium is precipitated as a white crystalline precipitate (CsCl.SbCl3), which is filtered off, and washed, when cold, with strong hydrochloric acid; since it is decomposed by water or on warming. The precipitate is washed into a beaker, and treated with sulphuretted hydrogen; after filtering off the sulphide of antimony, the solution leaves, on evaporation, the caesium as chloride.

Rubidium occurs widely diffused in nature, but in very small quantities. It is generally associated with caesium.

It is detected by the spectroscope, which shows two violet lines and two dark red ones. Like caesium, it is precipitated with platinic chloride, and in the ordinary course of work would be weighed as potassium. It is separated from potassium by fractional precipitation with platinic chloride. Rubidium platino-chloride is much less soluble than the potassium salt.

It is usual to look upon the salts of ammonia as containing a compound radical (NH4= Am), which resembles in many respects the metals of the alkalies. Ammonium occurs in nature as chloride in sal ammoniac (AmCl), as sulphate in mascagnine (Am2SO4), as phosphate in struvite (AmMgPO4.12H2O). Minerals containing ammonium are rare, and are chiefly found either in volcanic districts or associated with guano. Ammonia and ammonium sulphide occur in the waters of certain Tuscan lagoons, which are largely worked for the boracic acid they contain. The crude boracic acid from this source contains from 5 to 10 per cent. of ammonium salts. It is from these that the purer forms of ammonium compounds of commerce known as "from volcanic ammonia" are derived. But the bulk of the ammonia of commerce is prepared from the ammoniacal liquors obtained as bye-products in the working of certain forms of blast furnaces and coke ovens, and more especially in gas-making.

Ammonia hardly comes within the objects of assaying; but it is largely used in the laboratory, and the assayer is not unfrequently called on to determine it. Ammonium salts are mostly soluble in water. In strong solutions they give a yellow precipitate of ammonium platino-chloride on the addition of chloride of platinum; and with the acid tartrate of soda yield a white precipitate of hydric ammonic tartrate. These reactions are similar to those produced with potassium compounds.

Heated with a base, such as lime or sodic hydrate, ammonium salts are decomposed, yielding ammonia gas (NH3), which is readily soluble in water. The solution of this substance is known as ammonic hydrate or "ammonia."

They are volatilised on ignition; either with, or without, decomposition according to the acid present. This fact is of importance in analytical work; since it allows of the use of alkalinesolutions and reagents which leave nothing behind on heating. It must be remembered, however, that, although ammonic chloride is volatile, it cannot be volatilised in the presence of substances which form volatile chlorides without loss of the latter. For example: ferric oxide and alumina are thus lost, volatilising as chlorides; and there are some other compounds (notably ammonic magnesic arsenate) which on heating to redness suffer reduction. The presence of ammonic chloride in such cases must be avoided.

Detection.—Compounds of ammonium are detected by their evolving ammonia when mixed or heated with any of the stronger bases. The ammonia is recognised by its odour, by its alkaline reaction with litmus paper, and by yielding white fumes, when brought in contact with fuming acid. In consequence of the use of ammonium salts and ammonia as reagents, it is necessary to make a special test for and determination of ammonium.[94]In the ordinary course of work it will be "lost on ignition." The determination presents little difficulty, and is based on the method used for its detection.

Fig. 61.

Solution and Separation.—Although ammonium salts are soluble in water, there is no necessity for dissolving them. The compound containing the ammonia is boiled with an alkaline solution; and the liberated ammonia condensed and collected. The substance is weighed out into a flask of about 200 c.c. capacity. The flask is closed with a rubber cork perforated to carry a 20 c.c. pipette and a bulb exit tube. The latter is connected with a receiver, which is a small flask containing dilute hydrochloric acid (fig. 61). The flask containing the substance is corked, and the greater part of the soda solution is run in from the pipette. The solution is then boiled. The ammonia volatilises, and is carried over into the hydrochloric acid, with which it combines to form ammonic chloride. The distillation is carried on gently until the bulk of the liquid is driven over. The ammonia in the receiver will be mixed only with the excess of hydrochloric acid. This separation is used in all determinations.

The contents of the flask are transferred to a weighed platinum dish, and evaporated on the water-bath. It is dried until the weight is constant. The chloride of ammonium remains as a white mass which, after cooling in a desiccator, is weighed. It contains 33.72 per cent. of ammonium (NH4), or 31.85 per cent. of ammonia (NH3). On heating over the Bunsen burner it is completely volatilised, leaving no residue.

Weigh up 1.7 gram of the substance and place it in the flask. Measure off 50 c.c. of the normal solution of acid, place them in the receiver, and dilute with an equal volume of water. Run in through the pipette (by opening the clip) 20 c.c. of a strong solution of soda, boil until the ammonia has passed over, and then aspirate a current of air through the apparatus. Disconnect the receiver, and tint its contents with methyl orange. Titrate the residual acid with a semi-normal solution of alkali. Divide the c.c. of the "alkali" solution used by 2, and deduct from the 50 c.c. The difference will give the number of c.c. of the normal acid solution neutralised by the ammonia distilled over. Each c.c. of "acid" so neutralised, represents 1 per cent. of ammonia in the sample. If the results are to be reported as ammonium, 1.8 gram of the sample is taken instead of 1.7 gram.

This is effected by means of "Nessler's" reagent, which strikes a brown colour with traces of ammonia, even with a few hundredths of a milligram in 100 c.c. of liquid. With larger quantities of ammonia the reagent gives a precipitate. This reagent is a strongly alkaline solution of potassic mercuric iodide; and is thus made:—

Nessler's solution: Dissolve 17 grams of mercuric chloride in 300 c.c. of water; and add the solution to one of 35 grams of potassium iodide in 100 c.c. of water until a permanent precipitate is produced. Both solutions must be cold. Then make up to a litre by adding a 20 per cent. solution of potash. Add more of the mercuric chloride (a little at a time) until a permanent precipitate is again formed. Allow to settle, decant, and use the clear liquor. Four or five c.c. are used for each 100 c.c. of liquid to be tested.

A Standard Solution of Ammoniais made by dissolving 0.315gram of ammonic chloride in water, and diluting to 100 c.c. Ten c.c. of this are taken and diluted to 1 litre. One c.c. contains 0.01 milligram of ammonia (NH3).

In working, the solution containing the ammonia is diluted to a definite volume, and to such an extent that 50 c.c. of it shall not contain more than 0.02 or 0.03 milligram of ammonia. Fifty c.c. of it are transferred to a Nessler glass and mixed with 2 c.c. of Nessler's reagent. The colour is noted, and an estimate made as to the amount of ammonia it indicates. A measured quantity of the standard ammonia, judged to contain about as much ammonia as that in the assay, is then put into another Nessler glass. It is diluted to 50 c.c. with water, and mixed with 2 c.c. of "Nessler." After standing a minute or two, the colours in the two glasses are compared. If the tints are equal, the assay is finished; but if the standard is weaker or stronger than the assay, another standard, containing more or less ammonia, as the case may be, must be prepared and compared with the assay. Two such experiments will generally be sufficient; but, if not, a third must be made. The addition of more standard ammonia to the solution to which the "Nessler" has already been added does not give a satisfactory result.

When the ammonia in 50 c.c. has been determined, that in the whole solution is ascertained by a suitable multiplication. By 10, for example, if the bulk was 500 c.c., or by 20 if it was a litre.

Distilled water is used throughout. It must be free from ammonia; and is best prepared by distilling an ammonia-free spring water.

FOOTNOTES:[90]Al2Cl6+ 3Na2S2O3+ 3H2O = Al2(HO)6+ 6NaCl + 3S + 3SO2[91]3BeO,Al2O3,6SiO2[92]CaC2O4= CaCO3+CO.[93]Resolved into two with a powerful spectroscope.[94]Ammonium compounds are frequently produced when dissolving metals in nitric acid; or when nitrates are heated in the presence of the metals.

[90]Al2Cl6+ 3Na2S2O3+ 3H2O = Al2(HO)6+ 6NaCl + 3S + 3SO2

[91]3BeO,Al2O3,6SiO2

[92]CaC2O4= CaCO3+CO.

[93]Resolved into two with a powerful spectroscope.

[94]Ammonium compounds are frequently produced when dissolving metals in nitric acid; or when nitrates are heated in the presence of the metals.

Oxygen occurs in nature in the free state, forming 23 per cent. by weight, or 21 per cent. by volume of the atmosphere; but, since it is a gas, its presence is easily overlooked and its importance underestimated. Except in the examination of furnace-gases, &c., the assayer is not often called upon to determine its quantity, but it forms one of his most useful reagents, and there are many cases where he cannot afford to disregard its presence. It occurs not only in the air, but also dissolved in water; ordinary waters containing on an average 0.00085 per cent. by weight, or 0.85 parts per 100,000.

Chemically, it is characterised by its power of combining, especially at high temperatures, with the other elements, forming an important class of compounds called oxides. This combination, when rapid, is accompanied by the evolution of light and heat; hence oxygen is generally called the supporter of combustion. This property is taken advantage of in the operation of calcining, scorifying, cupelling, &c. The importance of a free access of air in all such work is seen when it is remembered that 1 litre of air contains 0.2975 gram of oxygen, and this quantity will only oxidise 0.1115 gram of carbon, 0.2975 gram of sulphur, or 3.849 grams of lead.

Oxidation takes place at the ordinary temperature with many substances. Examples of such action are seen in the weathering of pyrites, rusting of iron, and (in the assay office) the weakening of solutions of many reducing agents.

For methods of determining the percentage of oxygen in gases, for technical purposes, the student is referred to Winkler & Lunge's "Technical Gas Analysis."

Oxides are abundant in nature, almost all the commonly occurring bodies being oxidised. Water (H2O) contains 88.8 per cent. of oxygen; silica, lime, alumina, magnesia, and the other earths are oxides, and the oxides of the heavier metals are in many cases important ores; as, for example, cassiterite (SnO2), hæmatite (Fe2O3), magnetite (Fe3O4), and pyrolusite (MnO2). In fact, the last-named mineral owes its value to the excess of oxygen it contains, and may be regarded as an ore of oxygen rather than of manganese.

Most of the metals, when heated to redness in contact with air, lose their metallic lustre and become coated with, or (if the heating be prolonged) altogether converted into, oxide. This oxide was formerly termed a "calx," and has long been known to weigh more than the metal from which it was obtained. For example, one part by weight of tin becomes, on calcining, 1.271 parts of oxide (putty powder). The student will do well to try the following experiments:—Take 20 grams of tin and heat them in a muffle on a scorifier, scraping back the dross as it forms, and continuing the operation until the whole of the metal is burnt to a white powder and ceases to increase in weight.[95]Take care to avoid loss, and, when cold, weigh the oxide formed. The oxide should weigh 25.42 grams, which increase in weight is due to the oxygen absorbed from the air and combined with the metal. It can be calculated from this experiment (if there has been no loss) that oxide of tin contains 21.33 per cent. of oxygen and 78.67 per cent. of tin. Oxidation is performed with greater convenience by wet methods, using reagents, such as nitric acid, which contain a large proportion of oxygen loosely held. Such reagents are termed oxidising agents. Besides nitric acid, permanganate of potash, bichromate of potash, and peroxide of hydrogen are largely used for this purpose. One c.c. of nitric acid contains as much oxygen as 2.56 litres of air, and the greater part of this is available for oxidising purposes. Try the following experiment:—Take 2 grams of tin and cover in a weighed Berlin dish with 20 c.c. of dilute nitric acid, heat till decomposed, evaporate to dryness, ignite, and weigh. The 2 grams of tin should yield 2.542 grams of oxide. The increase in weight will be proportionally the same as in the previous experiment by calcination, and is due to oxygen, which in this case has been derived from the nitric acid.

The percentage of oxygen in this oxide of tin (or in any of the oxides of the heavier metals) may be directly determined by heating such oxides in a current of hydrogen, and collecting and weighing the water formed.

It is found by experiment that 88.86 parts by weight of oxygen, combining with 11.14 parts of hydrogen, form 100 parts of water; so that from the weight of water formed it is easy to calculate the amount of oxygen the oxide contained.

Fig. 62.

Take 1 gram of the dried and powdered oxide and place it in a warm dry combustion tube. Place the tube in a furnace, and connect at one end with a hydrogen apparatus provided with a sulphuric acid bulb for drying the gas, and at the other with a weighed sulphuric acid tube for collecting the water formed. The apparatus required is shown in fig. 62. Pass hydrogen through the apparatus, and, when the air has been cleared out, light the furnace. Continue the heat and current of hydrogen for half an hour (or longer, if necessary). Allow to cool. Draw a current of dry air through the weighed tube. Weigh. The increase in weight gives the amount of water formed, and this, multiplied by 0.8886, gives the weight of the oxygen. The percentage of oxygen thus determined should be compared with that got by the oxidation of the metal. It will be practically the same. The following results can be taken as examples:—

Twenty grams of tin, calcined as described, gave 25.37 grams of oxide.

Two grams of tin, oxidised with nitric acid and ignited, gave 2.551 grams of oxide.

One gram of the oxide of tin, on reduction in a current of hydrogen, gave 0.2360 gram of water (equivalent to 0.2098 gram of oxygen), and left 0.7900 gram of metal.

Ten grams of ferrous sulphate gave, on strong ignition, 2.898 grams of ferric oxide (Fe2O3)[96]instead of 2.877.

The student should similarly determine the percentage of oxygen in oxides of copper and iron. The former oxide may be prepared by dissolving 5 grams of copper in 50 c.c. of dilute nitricacid, evaporating to dryness, and strongly igniting the residue. The oxide of iron may be made by weighing up 10 grams of powdered ferrous sulphate (= to 2.014 grams of iron) and heating, at first gently, to drive off the water, and then at a red heat, until completely decomposed. The weight of oxide, in each case, should be determined; and the percentage of oxygen calculated. Compare the figures arrived at with those calculated from the formula of the oxides, CuO and Fe2O3.

It would be found in a more extended series of experiments that the same metal will, under certain conditions, form two or more oxides differing among themselves in the amount of oxygen they contain. These oxides are distinguished from one another by such names as "higher" and "lower oxides," "peroxides," "protoxides," "dioxides," &c.

The oxides may be conveniently classified under three heads:—

(1)Those that are reduced to metal by heat alone, such as the oxides of mercury, silver, platinum, gold, &c.;

(2)Those which are reduced by hydrogen at a red heat, which includes the oxides of the heavy metals;

(3)Those which are not reduced by these means, good examples of which are silica, alumina, the alkalies, and the alkaline earths.

Another important classification is into acid, basic and neutral oxides. The oxides of the non-metallic elements, such as sulphur, carbon, phosphorus, &c., are, as a rule, acid; and the more oxygen they contain, the more distinctly acid they are. The oxides of the metals are nearly all basic; and, as a rule, the less oxygen they contain, the more distinctly basic they are.

The basic oxides, which are soluble in acids, give rise to the formation of salts when dissolved therein. During the solution, water is formed, but no gas is evolved. The oxide dissolved in each case neutralizes an equivalent of the acid used for solution.[97]The basic properties of many of these can be taken advantage of for their determination. This is done in the case of soda, potash, lime, &c., by finding the quantity of acid required to neutralize a given weight of the substance.

There are some oxides which, under certain conditions, are acid to one substance (a stronger base) and basic to another (a stronger acid). For example, the oxides of lead and of tin, as also alumina, dissolve in caustic soda, acting as acids; whilst, onthe other hand, they combine with sulphuric or hydrochloric acid, playing the part of bases.

The oxides known as "earths," when ignited, are many of them insoluble in acids, although easily dissolved before ignition.

It is common in complete analyses of minerals to meet with cases in which the sum total of the elements found falls short of the amount of ore taken; and here oxygen must be looked for. For example, this occurs in the case of a mixture of pyrites with oxide of iron, or in a mixture of sulphides and sulphates. The state in which the elements are present, and the percentage (say of sulphides and sulphates) can in many cases be determined; but this is not always required. When the difference between the sum total and the elements found is small, it is reported as "oxygen and loss." When, however, it is considerable, the oxygen may be reported as such; and its amount be either determined directly in the way already described, or calculated from the best determination that can be made of the relative amounts of oxides, sulphides, sulphates, &c., present. Such cases require a careful qualitative analysis to find out that the substance is present; and then the separation of each constituent is made as strictly as possible. These remarks apply especially to ores of the heavy metals. The separation of the constituents is effected with suitable solvents applied in proper order. The soluble sulphates, for example, are extracted with water; the oxides by the dilute acids or alkalies in which they are known to be soluble. The oxygen in the sulphates and oxides thus obtained is estimated by determining the sulphur and metals in the solutions, and calculating the amount of oxygen with which they combine. The metals of the earths and alkalies are almost invariably present as oxides, and are reported as such; except it is known that they are present in some other form, such as fluoride or chloride. Thus, silica, alumina, lime, water, &c., appear in an analysis; even in those cases where "oxygen and loss" is also mentioned. As an example of such a report, take the following analysis of Spanish pyrites:—

Sulphur49.00Iron43.55Copper3.20Arsenic0.47Lead0.93Zinc0.35Lime0.10Silica, &c.0.63Water0.70Oxygen and loss1.07———100.00

The following example will illustrate the mode of calculating and reporting. A mineral, occurring as blue crystals soluble in water, and found on testing to be a mixed sulphate of iron and copper, gave on analysis the following results:—

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