OXYGEN AND OXIDES.—THE HALOGENS.

OXYGEN.

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.

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.

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 nitric acid, 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, on the 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:—

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

Water44.51per cent.
Sulphuric oxide28.82"
Copper8.44"
Ferrous iron11.81"
Ferric iron0.38"
Zinc0.28"
———
94.24

There is here a deficiency of 5.76 per cent. due to oxygen. Nothing else could be found, and it is known that in the sulphates the metals exist as oxides. By multiplying the weight of the copper by 1.252, the weight of copper oxide (CuO) will be ascertained; in this case it equals 10.57 per cent. The ferrous iron multiplied by 1.286 will give the ferrous oxide (FeO); in this case 15.19 per cent. The ferric iron multiplied by 1.428 will give the ferric oxide (Fe2O3); in this case 0.54 per cent. The zinc multiplied by 1.246 will give the zinc oxide (ZnO); in this case it equals 0.35 per cent. The analysis will be reported as—

Water44.51
Sulphuric oxide28.82
Copper oxide10.57equal to copper 8.44%
Ferrous oxide15.19
Ferric oxide0.54
Zinc oxide0.35
———
99.98

The following (A) is an analysis of a sample of South American copper ore, which will serve as a further illustration. The analysis showed the presence of 6.89 per cent. of ferrous oxide, and some oxide of copper.

The analysis (B) is that of an ore from the same mine after an imperfect roasting. It will be seen that the carbonates have been converted into sulphates. If the total sulphur simply had been determined, and the sulphate overlooked, the "oxygen and loss" would have been 5.65 per cent., an amount which would obviously require an explanation.

A.B.
Water0.250.59
Organic matter0.54—
Sulphur29.5021.33
Copper10.929.80{Copper9.57
{Copper oxide0.28
Iron32.0939.73{Iron34.32
{Ferric oxide7.73
Lead0.350.12
Zinc0.860.69
Cobalt0.060.11
Lime5.257.69
Magnesia2.332.55
Sulphuric oxide1.005.30
Carbon dioxide8.87—
"Insoluble silicates"5.128.38
Oxygen and loss2.862.47
———Potash0.15
100.00Soda1.09
———
100.00

WATER.

Water occurs in minerals in two forms, free and combined. The term "moisture" ought, strictly, to be limited to the first, although, as has already been explained, it is more convenient in assaying to apply the term to all water which is driven off on drying at 100° C. The combined water is really a part of the mineral itself, although it may be driven off at a high temperature, which varies with the base. In some cases a prolonged red heat is required; whilst with crystallised salts it is sometimes given off at the ordinary temperatures. This latter phenomenon, known as efflorescence, is mostly confined to artificial salts.

The determination of the combined water may often be made by simply igniting the substance from which the moisture has been removed. The quantity of water may be determined, either indirectly by the loss, or directly by collecting it in a calcium chloride tube, and weighing. In some cases, in which the loss on ignition does not give simply the proportion of combined water, it can be seen from the analysis to what else the loss is due; and, after a proper deduction, the amount of water can be estimated. For example, 1 gram of crystallised iron sulphate was found to contain on analysis 0.2877 gram of sulphuric oxide; and on igniting another gram, 0.2877 gram of ferric oxide was left. As the salt is known to be made up of ferrous oxide, sulphuric oxide, and combined water, the combined water can be thus calculated: 0.2877 gram of ferric oxide is equal to 0.2589 gram of ferrous oxide,[98] and consequently, the loss on ignition has been diminished by 0.0288 gram, which is the weight of oxygen absorbed by the ferrous oxide during calcining. The loss on ignition was 0.7123 gram, to which must be added 0.0288 gram; hence 0.7411 gram is the weight of the combined sulphuric oxide and water present. Deducting the weight of sulphuric oxide found, 0.2877 gram, there is left for combined water 0.4534 gram. The composition of 1 gram of the dry salt is then:—

Water0.4534
Sulphuric oxide0.2877
Ferrous oxide0.2589
———
1.0000

The following is another example:—A sample of malachite lost on ignition 28.47 per cent., leaving a residue which was found on analysis to be made up of oxide of copper (equal to 70.16 per cent. on the mineral), and silica and oxide of iron (equal to 1.37 per cent.). Carbon dioxide and water (but nothing else) was found to be present, and the carbon dioxide amounted to 19.64 per cent.; deducting this from the loss on ignition, we have 8.82 as the percentage of water present. The analysis was then reported as follows:—

Cupric oxide70.16equal to 56.0% copper.
Silica and ferric oxide1.37
Carbon dioxide19.64
Water8.82
———
99.99

Direct Determination of Combined Water.—Transfer about 3 grams of the substance to a piece of combustion tube (8 or 10 inches long), attached (as in fig. 63) at one end to a U-tube containing sulphuric acid, and at the other end to a calcium chloride tube. The last is weighed previous to the determination. The tube should be warmed to ensure complete dryness, and must be free from a misty appearance. Aspirate a current of air through the apparatus, heat the mineral by means of a Bunsen burner, cautiously at first, and afterwards to redness (if necessary). The water is driven off and condenses in the calcium chloride tube, which is afterwards cooled and weighed. The increase in weight is due to the water. If the substance gives off acid products on heating, it is previously mixed with some dry oxide of lead or pure calcined magnesia.

EXAMINATION OF WATERS.

The assayer is occasionally called on to test water for the purpose of ascertaining the nature and quantity of the salts contained in it, and whether it is or is not fit for technical and drinking purposes.

In mineral districts the water is generally of exceptional character, being more or less charged, not only with earthy salts, but also frequently with those of the metals. Distilled water is only used by assayers in certain exceptional cases, so that by many it would be classed among the rarer oxides. Water of ordinary purity will do for most purposes, but the nature and quantity of the impurities must be known.

The following determinations are of chief importance:—

Total Solids at 100° C.—Where simply the amount is required, take 100 c.c. and evaporate on the water-bath in a weighed dish; then dry in the water-oven, and weigh.

Total Solids Ignited.—The above residue is very gently ignited (keeping the heat well below redness), and again weighed. A larger loss than 4 or 5 parts per 100,000 on the water requires an explanation.

Chlorine.—Take 100 c.c. of the water in a porcelain dish, add 2 c.c. of a 5 per cent. solution of neutral potassic chromate, and titrate with a neutral standard solution of nitrate of silver, made by dissolving 4.789 grams of crystallised silver nitrate in distilled water, and diluting to 1 litre. The addition of the nitrate of silver is continued until the yellow of the solution assumes a reddish tint. The reaction is very sharp. Each c.c. of nitrate of silver used is equal to 1 part by weight of chlorine in 100,000 of water. At inland places this rarely amounts to more than 1 in 100,000; but near the sea it may amount to 3 or 5. More than this requires explanation, and generally indicates sewage pollution.

Nitric Pentoxide (N2O5).—It is more generally reported under the heading, "nitrogen as nitrates." Take 250 c.c. of the water and evaporate to 2 or 3 c.c.; acidulate with a few drops of dilute sulphuric acid, and transfer to a nitrometer (using strong sulphuric acid to wash in the last traces). The sulphuric acid must be added to at least twice the bulk of the liquid. Shake up with mercury. The mercury rapidly flours, and nitric oxide is given off (if any nitrate is present). The volume of the nitric oxide (corrected to normal temperature and pressure), multiplied by 0.25, gives the parts of nitrogen per 100,000; or, multiplied by 0.965, will give the nitric pentoxide in parts per 100,000. In well and spring waters the nitrogen may amount to 0.3 or 0.4 parts per 100,000; or in richly cultivated districts 0.7 or 0.8 parts per 100,000. An excess of nitrates is a suspicious feature, and is generally due to previous contamination.

Ammonia.—Take 500 c.c. of the water and place them in a retort connected with a Liebig's condenser. Add a drop or two of a solution of carbonate of soda and distil over 100 c.c.; collect another 50 c.c. separately. Determine the ammonia in the distillate colorimetrically (with Nessler's solution, as described under Ammonia) and compare with a standard solution of ammonic chloride containing 0.0315 gram of ammonic chloride in 1 litre of water. One c.c. contains 0.01 milligram of ammonia. The second distillate will show little, if any, ammonia in ordinary cases. The amounts found in both distillates are added together, and expressed in parts per 100,000.

Waters (other than rain and tank waters) which contain more than 0.003 per 100,000 are suspicious.

Organic Matter.—The organic matter cannot be determined directly; but for ordinary purposes it may be measured by the amount of permanganate of potassium which it reduces, or by the amount of ammonia which it evolves on boiling with an alkaline permanganate of potassium solution.

A. Albuminoid Ammonia.—To the residue left after distilling the ammonia add 50 c.c. of a solution made by dissolving 200 grams of potash and 8 grams of potassium permanganate in 1100 c.c. of water, and rapidly boiling till the volume is reduced to 1 litre (this should be kept in a well stoppered bottle, and be occasionally tested to see that it is free from ammonia). Continue the distillation, collecting 50 c.c. at a time, until the distillate is free from ammonia. Three or four fractions are generally sufficient. Determine the ammonia colorimetrically as before. If the total albuminoid ammonia does not exceed 0.005 in 100,000, the water may be regarded as clean as regards organic matter; if it amounts to more than 0.015, it is dirty.

B. Oxygen Consumed.—A standard solution of permanganate of potash is made by dissolving 0.395 gram of the salt in water and diluting to 1 litre. Each c.c. equals 0.1 milligram of available oxygen. The following are also required:—1. A solution of sodium hyposulphite containing 1 gram of the salt (Na2S2O3.5H2O) in 1 litre of water. 2. Dilute sulphuric acid, made by adding one part of the acid to three of water, and titrating with the permanganate solution till a faint pink persists after warming for several hours. 3. Starch paste. 4. Potassium iodide solution.

Take 250 c.c. of the water in a stoppered bottle, add 10 c.c. of sulphuric acid and 10 c.c. of the permanganate, and allow to stand in a warm place for four hours. Then add a few drops of the solution of potassium iodide, and titrate the liberated iodine with "hypo," using starch paste towards the end as an indicator. To standardise the hyposulphite, take 250 c.c. of water and 10 c.c. of sulphuric acid, and a few drops of potassium iodide; then run in 10 c.c. of the "permanganate" solution, and again titrate; about 30 c.c. of the "hypo" will be used. The difference in the two titrations, divided by the last and multiplied by 10, will give the c.c. of permanganate solution used in oxidising the organic matter in the 250 c.c. of water. Each c.c. represents 0.04 parts of oxygen in 100,000.

Metals.—These may for the most part be estimated colorimetrically.

Lead.—Take 100 c.c. of the water in a Nessler tube, and add 10 c.c. of sulphuretted hydrogen water, and compare the tint, if any, against a standard lead solution, as described under Colorimetric Lead. Report in parts per 100,000.

Copper.—Proceed as with the last-mentioned metal; but, if lead is also present, boil down 500 c.c. to about 50 c.c., then add ammonia, filter, and estimate the copper in the blue solution, as described under Colorimetric Copper.

Iron.—Take 50 c.c., or a smaller quantity (if necessary), dilute up to the mark with distilled water, and determine with potassium sulphocyanate, as described under Colorimetric Iron.

Zinc.—Zinc is the only other metal likely to be present; and, since it cannot be determined colorimetrically, it must be separately estimated during the examination of the "total solids."

Examination of "Total Solids."—Evaporate 500 c.c. to dryness with a drop or two of hydrochloric acid. Take up with hydrochloric acid, filter, ignite, and weigh the residue as "silica." To the filtrate add a little ammonic chloride and ammonia, boil and filter, ignite, and weigh the precipitate as "oxide of iron and alumina." Collect the filtrate in a small flask, add a few drops of ammonium sulphide or pass sulphuretted hydrogen, cork the flask, and allow to stand overnight; filter, wash, and determine the zinc gravimetrically as oxide of zinc. If copper or lead were present, they should have been previously removed with sulphuretted hydrogen in the acid solution. To the filtrate add ammonic oxalate and ammonia, boil for some time, allow to stand, filter, wash, ignite, and weigh as "lime." Evaporate the filtrate with nitric acid, and ignite. Take up with a few drops of dilute hydrochloric acid, add baric hydrate in excess, evaporate, and extract with water. The residue contains the magnesia; boil with dilute sulphuric acid, filter, precipitate it with phosphate of soda and ammonia, and weigh as pyrophosphate. The aqueous extract contains the alkalies with the excess of barium. Add sulphuric acid in slight excess, filter, evaporate, and ignite strongly. The residue consists of the sulphates of the alkalies (which are separately determined, as described under Potash).

Sulphuric Oxide (SO3).—Take 200 c.c. and boil to a small bulk with a little hydrochloric acid, filter (if necessary), add baric chloride solution in slight excess to the hot solution, filter, ignite, and weigh as baric sulphate.

Carbon Dioxide (free).—Carbon dioxide exists in waters in two forms, free and combined. The latter generally occurs as bicarbonate, although on analysis it is more convenient to consider it as carbonate, and to count the excess of carbon dioxide with the free. The method is as follows:—To determine the free carbon dioxide, take 100 c.c. of the water, place them in a flask with 3 c.c. of a strong solution of calcium chloride and 2 c.c. of a solution of ammonic chloride, next add 50 c.c. of lime-water. The strength of the lime-water must be known. Make up to 200 c.c. with distilled water, stop the flask, and allow the precipitate to settle. Take out 100 c.c. of the clear solution with a pipette, and titrate with the standard solution of acid.[99] The number of c.c. required, multiplied by two, and deducted from that required for the 50 c.c. of lime-water, and then multiplied by 0.0045, will give the carbon dioxide present other than as normal carbonates.

Carbon Dioxide combined as normal carbonate.—100 c.c. of the water are tinted with phenacetolin or lacmoid; then heated to near boiling, and titrated with standard acid. The number of c.c. used, multiplied by 0.0045, will give the weight in grams of the combined carbon dioxide.

Free Acid.—In some waters (especially those from mining districts) there will be no carbonates. On the contrary, there may be free mineral acid or acid salts. In these cases it is necessary to determine the amount of acid (other than carbon dioxide) present in excess of that required to form normal salts. This is done in the following way:—Make an ammoniacal copper solution by taking 13 grams of copper sulphate (CuSO4.5H2O), dissolving in water, adding solution of ammonia until the precipitate first formed has nearly dissolved, and diluting to 1 litre. Allow to settle, and decant off the clear liquid. The strength of this solution is determined by titrating against 10 or 20 c.c. of the standard solution of sulphuric acid (100 c.c. = 1 gram H2SO4). The finishing point is reached as soon as the solution becomes turbid from precipitated cupric hydrate. At first, as each drop falls into the acid solution, the ammonia and cupric hydrate combine with the free acid to form ammonic and cupric sulphates; but as soon as the free acid is used up, the ammonia in the next drop not only precipitates an equivalent of cupric hydrate from the solution, but also throws down that carried by itself. This method is applicable in the presence of metallic sulphates other than ferric. The standardising and titration should be made under the same conditions. Since sulphuric acid and sulphates are predominant in waters of this kind, it is most convenient to report the acidity of the water as equivalent to so much sulphuric acid.

Dissolved Oxygen.—For the gasometric method of analysing for dissolved oxygen, and for the Schützenberger's volumetric method, the student is referred to Sutton's "Volumetric Analysis." The following is an easy method of estimating the free oxygen in a water:—Take 20 c.c. of a stannous chloride solution (about 20 grams of the salt with 10 c.c. of hydrochloric acid to the litre); add 10 c.c. of hydrochloric acid, and titrate in an atmosphere of carbon dioxide with standard permanganate of potassium solution (made by dissolving 1.975 gram of the salt in 1 litre of water: 1 c.c. equals 0.5 milligram of oxygen). A similar titration is made with the addition of 100 c.c. of the water to be tested. Less permanganate will be required in the second titration, according to the amount of oxygen in the water; and the difference, multiplied by 0.5, will give the weight of the oxygen in milligrams. Small quantities of nitrates do not interfere.

In reporting the results of the analysis, it is customary to combine the acids and bases found on some such principle as the following:—The sulphuric oxide is calculated as combined with the potash, and reported as potassic sulphate (K2SO4); the balance of the sulphuric oxide is then apportioned to the soda, and reported as sulphate of soda (Na2SO4); if any is still left, it is reported as calcium sulphate (CaSO4), and after that as magnesic sulphate (MgSO4). When the sulphuric oxide has been satisfied, the chlorine is distributed, taking the bases in the same order, then the nitric pentoxide, and lastly the carbon dioxide. But any method for thus combining the bases and acids must be arbitrary and inaccurate. It is extremely improbable that any simple statement can represent the manner in which the bases and acids are distributed whilst in solution; and, since different chemists are not agreed as to any one system, it is better to give up the attempt, and simply state the results of the analysis. This has only one inconvenience. The bases are represented as oxides; and, since some of them are present as chlorides, the sum total of the analysis will be in excess of the actual amount present by the weight of the oxygen equivalent to the chlorine present as chloride. The following is an example of such a statement:—

Parts per 100,000.
Total solids, dried at 100° C.28.73
Chlorine1.70
Nitrogen as nitrate0.03
Ammonia0.001
Albuminoid ammonia0.004
"Oxygen consumed" in 4 hours0.01

The solids were made up as under:—

Per 100,000 of the Water.
Potash0.38
Soda2.01
Magnesia1.44
Lime10.55
Ferric oxide0.01
Silica0.30
Sulphuric oxide3.69
Nitrogen pentoxide0.11
Carbon dioxide8.38
Chlorine1.70
Volatile and organic matter0.66
———
29.23
Less oxygen equivalent to chlorine found0.39
———
28.84

For the preparation of distilled water, the apparatus shown in fig. 64 is convenient for laboratory use. It consists of a copper retort heated by a ring gas-burner, and connected with a worm-condenser.

PRACTICAL EXERCISE.

A mineral, on analysis, gave the following results:—Water, 44.94 per cent.; sulphuric oxide, 28.72 per cent.; ferrous iron, 13.92 per cent.; ferric iron, 0.35 per cent.; copper, 6.1 per cent. The mineral was soluble in water, and showed nothing else on testing. How would you report the analysis? Calculate the formula for the salt.

THE HALOGENS.

There is a group of closely allied elements to which the name halogen (salt-producer) has been given. It comprises chlorine, bromine, iodine, and fluorine. These elements combine directly with metals, forming as many series of salts (chlorides, bromides, iodides, and fluorides), corresponding to the respective oxides, but differing in their formulæ by having two atoms of the halogen in the place of one atom of oxygen. For example, ferrous oxide is FeO and ferrous chloride is FeCl2, and, again, ferric oxide is Fe2O3, whilst ferric chloride is Fe2Cl6. These salts differ from the carbonates, nitrates, &c., in containing no oxygen. Consequently, it is incorrect to speak of such compounds as chloride of potash, fluoride of lime, &c., since potash and lime are oxides. It is important to bear this in mind in reporting analyses in which determinations have been made, say, of chlorine, magnesia, and potash, or of fluorine, silica, and alumina. It is necessary in all such cases to deduct from the total an amount of oxygen equivalent to the halogen found, except, of course, where the base has been determined and recorded as metal. Compounds containing oxides and fluorides, &c., do not lend themselves to the method of determining the halogen by difference. For example, topaz, which, according to Dana, has the formula Al2SiO4F2, would yield in the ordinary course of analysis—

Alumina55.4%
Silica32.6
Fluorine20.6
———
108.6

The oxygen equivalent to 20.6 per cent. fluorine may be found by multiplying the percentage of fluorine by 0.421; it is 8.7 per cent., and must be deducted. The analysis would then be reported thus:—

Alumina55.4%
Silica32.6
Fluorine20.6
———
108.6
Less oxygen equivalent to fluorine8.7
———
99.9

Take as an illustration the following actual analysis by F.W. Clarke and J.S. Diller:—

Alumina57.38%
Silica31.92
Fluorine16.99
Potash0.15
Soda1.33
Water0.20
———
107.97
Deduct oxygen equivalent7.16
———
100.81

In calculating the factor for the "oxygen equivalent," divide the weight of one atom of oxygen (16) by the weight of two atoms of the halogen; for example, with chlorine it would be 16/71 or 0.2253; with bromine, 16/160 or 0.1000; with iodine, 16/254 or 0.063; and with fluorine, 16/38 or 0.421.

CHLORINE AND CHLORIDES.

Chlorine occurs in nature chiefly combined with sodium, as halite or rock salt (NaCl). With potassium it forms sylvine (KCl), and, together with magnesium, carnallite (KCl.MgCl2.6H2O). Of the metalliferous minerals containing chlorine, kerargyrite, or horn silver (AgCl), and atacamite, an oxychloride of copper (CuCl2.3Cu(HO)2.) are the most important. Apatite (phosphate of lime) and pyromorphite (phosphate of lead) contain a considerable amount of it. Chlorine is a gas of a greenish colour, possessing a characteristic odour, and moderately soluble in water. It does not occur native, and is generally prepared by the action of an oxidising agent on hydrochloric acid. It combines directly with metals at the ordinary temperature (even with platinum and gold), forming chlorides, which (except in the case of silver) are soluble.

It is important in metallurgy, because of the extensive use of it in extracting gold by "chloridising" processes. It is also used in refining gold.

Detection.—Compounds containing the oxides of chlorine are not found in nature, because of the readiness with which they lose oxygen. By reduction they yield a chloride; the form in which chlorine is met with in minerals. In testing, the compound supposed to contain a chloride is boiled with water, or, in some cases, dilute nitric acid. To the clear solution containing nitric acid a few drops of nitrate of silver solution are added. If, on shaking, a white curdy precipitate, soluble in ammonia, separates out, it is sufficiently satisfactory evidence of the presence of chlorides.

Solution and Separation.—The chlorides are generally soluble in water, and are got into solution by extracting with warm dilute nitric acid. Or, if insoluble, the substance is fused with carbonate of soda, extracted with water, and the filtrate acidified with nitric acid. For the determination, it is not necessary to obtain the solution of the chloride free from other acids or metals. If tin, antimony, mercury, or platinum is present, it is best to separate by means of sulphuretted hydrogen. The chloride is determined in the solution after removal of the excess of the gas. Where traces of chlorides are being looked for, a blank experiment is made to determine the quantity introduced with the reagents. One hundred c.c. of ordinary water contains from 1 to 3 milligrams of chlorine. On the addition of nitrate of silver to the nitric acid solution, chloride of silver separates out. This is free from other substances, except, perhaps, bromide and iodide.

GRAVIMETRIC DETERMINATION.

Freely mix the solution containing the chloride with dilute nitric acid, filter (if necessary), and treat with nitrate of silver. Heat nearly to boiling, and, when the precipitate has settled, filter, and wash with hot distilled water. Dry, and transfer to a weighed Berlin crucible. Burn the filter-paper separately, and convert any reduced silver into chloride by alternate treatment with drops of nitric and of hydrochloric acid. Add the main portion to this, and heat cautiously till the edges of the mass show signs of fusing (about 260°). Cool in the desiccator and weigh. The substance is chloride of silver (AgCl), and contains 24.73 per cent. of chlorine.

The precipitated chloride is filtered and washed as soon as possible after settling, since on exposure to light it becomes purple, and loses a small amount of chlorine.

VOLUMETRIC METHOD.

There are several volumetric methods; but that based on the precipitation of silver chloride in neutral solution, by means of a standard solution of silver nitrate (using potassium chromate as indicator), is preferred. Silver chromate is a red-coloured salt; and, when silver nitrate is added to a solution containing both chloride and chromate, the development of the red colour marks off sharply the point at which the chloride is used up. Silver chromate is decomposed and consequently decolorised by solution of any chloride. The solution for this method must be neutral, since free acid prevents the formation of the red silver chromate. If not already neutral, it is neutralised by titrating cautiously with a solution of soda. In a neutral solution, other substances (such as phosphates and arsenates) also yield a precipitate with a solution of nitrate of silver; and will count as chloride if they are not removed.

The Standard Solution of Nitrate of Silver is made by dissolving 23.94 grams of the salt (AgNO3) in distilled water, and diluting to 1 litre; 100 c.c. are equal to 0.5 gram of chlorine.

The indicator is made by adding silver nitrate to a strong neutral solution of yellow chromate of potash (K2CrO4), till a permanent red precipitate is formed. The solution is allowed to settle, and the clear liquid decanted into a stoppered bottle labelled "chromate indicator for chlorine."

Standardise the silver nitrate by weighing up 0.5 gram of pure sodium chloride (or potassium chloride). Transfer to a flask and dissolve in distilled water; dilute to 100 c.c. Fill an ordinary burette with the standard silver solution, and (after adjusting) run into the flask a quantity sufficient to throw down the greater part of the chlorine. Add a few drops of the chromate indicator and continue the addition of the silver nitrate until the yellow colour of the solution becomes permanently tinted red, after shaking. This shows that the chlorine is all precipitated, and that the chromate is beginning to come down. The further addition of a couple of drops of the silver solution will cause a marked difference in the tint. Read off the quantity run in, and calculate the standard. One gram of sodium chloride contains 0.6062 gram of chlorine; and 1 gram of potassium chloride contains 0.4754 gram.

For the determination of small quantities of chloride (a few milligrams), the same method is used; but the standard solution is diluted so that each c.c. is equal to 1 milligram of chlorine; and the chromate indicator is added before titrating. The standard solution is made by measuring off 200 c.c. of the solution described above, and diluting with distilled water to 1 litre.

BROMINE AND BROMIDES.

Bromine closely resembles chlorine in the nature of its compounds. It does not occur free in nature, but is occasionally found in combination with silver as bromargyrite (AgBr) and, together with chloride, in embolite. It mainly occurs as alkaline bromides in certain natural waters. Nearly all the bromine of commerce is derived from the mother liquors of salt-works—i.e., the liquors from which the common salt has been crystallised out. Bromine combines directly with the metals, forming a series of salts—the bromides. In ordinary work they are separated with, and (except when specially tested for) counted as, chlorides. They are detected by adding chlorine water to the suspected solution and shaking up with carbon bisulphide. Bromine colours the latter brown.

IODINE AND IODIDES.

Iodine does not occur in nature in the free state; and iodides are rare, iodargyrite or iodide of silver (AgI) being the only one which ranks as a mineral species. Iodates are found associated with Chili saltpetre, which is an important source of the element.

Iodine and Iodides are largely used in the laboratory, and have already been frequently referred to. It is used as an oxidising agent in a similar manner as permanganate and bichromate of potash, especially in the determinations of copper, arsenic, antimony, and manganese.

Iodine is not readily soluble in water; but dissolves easily in a concentrated solution of potassium iodide. Its solutions are strongly coloured; a drop of a dilute solution colours a large volume of water decidedly yellow; on the addition of starch paste, this becomes blue. The delicacy of this reaction is taken advantage of in titrations to determine when free iodine is present. The blue colour may be alternately developed and removed by the addition of iodine (or an oxidising agent) and hyposulphite of soda (or some other reducing agent). In decolorising, the solution changes from blue or black to colourless or pale yellow according to circumstances. Sometimes the solution, instead of remaining colourless, gradually develops a blue which recurs in spite of the further addition of the reducing agent. In these cases the conditions of the assay have been departed from, or (and this is more often the case) there is some substance present capable of liberating iodine.

Iodine forms a series of salts—the iodides—resembling in many respects the chlorides. These can be obtained by direct combination of the metals with iodine.

Detection.—Free iodine is best recognised by the violet vapours evolved from the solution on heating, and by the blue or black colour which it strikes on the addition of starch paste. Iodides are detected by boiling with strong solutions of ferric sulphate or chloride. Iodine is liberated, distilled over, and collected. Chlorine also liberates iodine from iodides; and this reaction is frequently made use of in assaying. A process based on this is described under Manganese. All substances which liberate chlorine on boiling with hydrochloric acid (dioxides, bichromates, permanganates, &c.) are determined in a similar way.

Solution and Separation.—Most iodides are soluble in water or dilute acids. The separation is effected by distilling the substance with solution of ferric sulphate, and collecting the vapour in a dilute solution of sulphurous acid or arsenite of soda. On the completion of the distillation, the iodine will be in the distillate as iodide; and the gravimetric determination is made on this.

GRAVIMETRIC DETERMINATION.

To the solution containing the iodine, as iodide, and which is free from chlorides (and bromides), add a little dilute nitric acid and nitrate of silver till no further precipitate is produced. Filter off, wash with hot water, and dry. Clean the filter-paper as much as possible, and burn it. Collect the ash in a weighed porcelain crucible, add the main portion, and heat to incipient fusion; cool, and weigh. The substance is silver iodide, and contains 54.03 per cent. of iodine.

VOLUMETRIC METHOD.

This is for the titration of free iodine, and is practically that which is described under Manganese. The substance to be determined is distilled with ferric sulphate, and the iodine is collected in a solution of potassium iodide, in which it readily dissolves. If flaky crystals separate out in the receiver, more potassium iodide crystals are added. When the distillation is finished, the receiver is disconnected, and its contents washed out into a beaker and titrated with "hypo." The standard solution of "hypo" is made by dissolving 19.58 grams of hyposulphite of soda (Na2S2O3.5H2O) in water and diluting to 1 litre; 100 c.c. are equal to 1 gram of iodine. To standardise the solution, weigh up 0.25 gram of pure iodine in a small beaker. Add 2 or 3 crystals of potassium iodide; cover with water; and, when dissolved, dilute to 50 or 100 c.c. Titrate, and calculate the standard.

FLUORINE AND FLUORIDES.

Fluorine is frequently met with as calcium fluoride or fluor-spar (CaF2). It occurs less abundantly as cryolite (Na3AlF6), a fluoride of aluminium and sodium, which is used in glass-making. Certain other rarer fluorides are occasionally met with. Fluorine is also found in apatite, and in some silicates, such as topaz, tourmaline, micas, &c.

Hydrofluoric acid is used for etching glass and opening up silicates. It attacks silica, forming fluoride of silicon (SiF4), which is volatile. Silica is by this means eliminated from other oxides, which, in the presence of sulphuric acid, are fixed. The commercial acid is seldom pure, and generally weak; and the acid itself is dangerously obnoxious. The use of ammonium fluoride (or sodium fluoride) and a mineral acid is more convenient. Determinations of this kind are made in platinum dishes enclosed in lead or copper vessels in a well-ventilated place. Fluor-spar is useful as a flux in dry assaying; it renders slags, which would otherwise be pasty, quite fluid. Fluorides generally are fusible, and impart fusibility to substances with which they form weak compounds. Their fluxing action does not depend on the removal of silicon as fluoride.

Detection.—Fluorides in small quantity are easily overlooked unless specially sought for. In larger amounts they are recognised by the property hydrofluoric acid has of etching glass. A watch-glass is warmed, and a layer of wax is melted over the convex side. When cold, some lines are engraved on the waxed surface with any sharp-pointed instrument. The substance to be tested is powdered; and moistened, in a platinum dish, with sulphuric acid. The watch-glass is filled with cold water and supported over the dish. The dish is then carefully warmed, but not sufficiently to melt the wax. After a minute or two, the glass is taken off, and the wax removed. If the substance contained fluorine, the characters will be found permanently etched on the glass. An equally good, but more rapid, test is to mix the powdered substance with some silica, and to heat the mixture in a test tube with sulphuric acid. Silicon fluoride is evolved, and, if a moistened glass rod is held in the tube, it becomes coated with a white deposit of silica, formed by the decomposition of the silicon fluoride by the water. This is also used as a test for silica; but in this case the substance is mixed with a fluoride, and the experiment must obviously be carried out in a platinum vessel.

Separation and Determination.—The determination of fluorine is difficult. In the case of fluorides free from silicates (such as fluor-spar), it is determined indirectly by decomposing a weighed portion with sulphuric acid, evaporating, igniting, and weighing the residual sulphate. The increase in weight multiplied by 0.655 gives the weight of fluorine.

In the presence of silica this method does not answer, because of the volatilisation of silicon fluoride. In these cases Wöhler adopted the following plan, which resembles that for the indirect determination of carbon dioxide. Mix the weighed substance in a small flask with powdered silica and sulphuric acid. The mouth of the flask is closed with a cork carrying a tube which is filled, the first half with calcium chloride and the second half with pumice coated with dried copper sulphate. The apparatus is weighed quickly, and then warmed till decomposition is complete. A current of dry air is aspirated for a minute or two; and the apparatus again weighed. The loss in weight gives that of the silicon fluoride (SiF4), which, multiplied by 0.7307, gives the weight of fluorine.

Fresenius uses the same reaction; but collects and weighs the silicon fluoride. The finely powdered and dried substance is mixed with ten or fifteen times its weight of ignited and powdered silica. The mixture is introduced into a small dry flask connected on one side with a series of drying-tubes, and on the other with an empty tube (to condense any sulphuric acid). To this last is joined a drying-tube containing chloride of calcium and anhydrous copper sulphate. This is directly connected with a series of three weighed tubes in which the fluoride of silicon is collected. The last of these is joined to another drying-tube. The first weighed tube contains pumice and cotton wool, moistened with water; the second tube contains soda-lime as well as (in the upper half of the second limb) fused calcium chloride between plugs of wool; the third tube is filled half with soda-lime and half with fused calcium chloride. The distilling-flask containing the substance mixed with silica is charged with 40 or 50 c.c. of sulphuric acid, and placed on the hot plate. Alongside it is placed a similar dry flask containing a thermometer, and the temperature in this is kept at 150° or 160° C. A current of air is sent through the tubes during the operation, which takes from one to three hours for from 0.1 to 1 gram of the substance. A correction is made by deducting 0.001 gram for every hour the dried air has been passed through. The increase in weight of the three tubes gives the weight of the silicon fluoride.

Penfield uses a similar arrangement, but passes his silicon fluoride into an alcoholic solution of potassium chloride. Silica and potassium silico-fluoride are precipitated, and hydrochloric acid is set free.[100] The acid thus liberated is titrated, with a standard solution of alkali, in the alcoholic solution, and from the amount of free acid found the fluorine is calculated. The weight of hydrochloric acid (HCl) found, multiplied by 1.562, gives the weight of the fluorine. With this method of working, fewer U-tubes are required. The exit tube from the flask is bent so as to form a small V, which is kept cool in water; this is directly connected with the U-tube containing the alcoholic solution of potassium chloride. The flask with the assay is heated for about two hours, and a current of dry air is aspirated throughout the determination. Fluoride of silicon is a gas not easily condensed to a liquid: but is immediately decomposed by water or moist air.