When acted on by heat and dilute acids the following anhydrides are produced—the formulæ given must be taken as to some extent provisional.

The white deposit which occurs on pit sides and in the interior of leather where gambier is largely used, and which is sometimes called "the whites," consists of catechin. It may be decomposed by warm sulphuric acid. This deposit is favoured by the use of hot gambier liquors. It is probable that by exposure to the air and by boiling, catechin is gradually converted into catechutannic acid during the tanning process, and hence the practical tanning value of cube gambier is probably greater than analysis indicates. Hunt has, however, shown (Jour. Soc. Chem. Ind., iv. 266) that, as estimated by the Löwenthal process, the tannin is to some extent lessened by boiling.

Kinoin, C₁₄H₁₂O₆ (Etti, Berl. Ber., xi. 1879), obtained from green or malabar kino, a product very similar to cutch, by boiling with dilute hydrochloric acid and extraction by agitation with ether, is very similar in its properties to catechin. It does not itself precipitate gelatin, but like catechin, yields a series of anhydrides or reds, which do so. On dry distillation it yields catechol and common phenol; and when heated with hydrochloric acid at 248°-266° F. (120°-130° C.) methyl chloride, catechol and gallic acid. Hence its constitution is probably that of methyl-catechol gallate.

Quebracho-catechin was found by P. N. Arata (Chem. Soc. Jour., xl. 1152) in the wood of quebracho colorado (p. 40), but in too small quantity for detailed investigation. It probably bears a similar relation to quebrachitannic acid that ordinary catechin does to catechutannic acid. It is insoluble in cold and only slightly soluble in hot water, but very soluble in alcohol and ether. Its solution is clouded by normal lead acetate, and gives rose-coloured precipitates with basic lead acetate and mercurous nitrate, and blackish with ferric acetate; it reduces silver-nitrate and gold chloride, and is coloured yellow by nitric acid, red by sulphuric acid, yellowish by sodium hypochlorite, and green by Fehling's solution. It does not precipitate gelatin, or alkaloids.

Catechutannic acid has been pretty fully described under catechin, of which it is the first anhydride (p. 80). It is possibly identical with mimo-tannic acid, the tannin of cutch and mimosa bark, which is chemically very similar, but greatly differs in its practical effect in tanning. For some further reactions of cutch and gambier infusions see p. 113. It gives a greyish-green precipitate with ferric salts, and (distinction from gallotannic acid) precipitates cupric sulphate but not tartar emetic.

Quebrachitannic acid is got from the wood of the quebracho colorado, Quebrachia lorentzii (formerly Loxopterygium) which must not be confounded with the bark of the Aspidospermum quebrachia, which is valuable, not for its tannin, but for an alkaloid, aspidospermin, which is used for medical purposes. It has been pretty thoroughly investigated by P. N. Arata (Chem. Soc. Jour., xxxiv. 986 and xl. 1152). It seems, however, a little doubtful to the writer whether the substance investigated by Arata was not the anhydride of the tannin, rather than the tannin itself, as it presents many points of analogy to catechu-red and was less soluble in water than quebracho tannin appears in practice to be.

According to Arata, quebrachitannic acid is a pale red amorphous mass, having an astringent taste and yielding a light cinnamon coloured powder. It is insoluble in carbon-bisulphide, turpentine oil, and benzene. Its aqueous solution gives a white precipitate with both normal and basic lead acetate, which when heated, acquires first a rose, and then a chocolate colour; with ferric chloride a green liquid is produced, which changes after a time to red and becomes black on addition of sodium acetate. It forms white precipitates with gelatin, albumen, and alkaloids. By dry distillation it yields catechol. By fusion with potash or the action of sulphuric acid, phloroglucol and protocatechuic acid, while nitric acid converts it into oxalic and picric acids. While it shows great similarity in its reactions to catechutannic acid, it differs materially in percentage composition, containing only 52·5 as compared with 62·0 per cent. of carbon.

CHAPTER V.

WATER AS USED IN TANNING.

Water, as obtained from rivers, wells, or water companies, contains a variety of impurities which affect its use in tanning, but of which in most cases the precise influence is very imperfectly known. These may be classified into (1) merely suspended matters, such as clay and mud, and sometimes animal or vegetable organisms such as infusoria; (2) dissolved mineral matters, which consist mostly of lime and magnesia salts and which make the water hard; (3) and organic dissolved impurities, such as the brown colour of peat water and the putrefying animal matters of sewage contamination.

Mud is always objectionable. It frequently contains organic slime and organisms which encourage the putrefaction of hides put in it to wash or soften. It also almost invariably contains iron as one of its constituents, and hence stains leather, and gives bad coloured liquors. It is not easily got rid of by filtration, as large filter-beds are expensive and difficult to keep in order, and much space is required to clear water by subsidence. Some filter easily cleaned offers the best chance of success. The Pulsometer Company supply such a filter, consisting of sponge tightly packed below a perforated piston. To cleanse the filter a stream of water is passed the reverse way, and the piston raised, and worked up and down, either by hand or power, so as to loosen and knead the sponge. The Atkins "water scrubber," in which sand may be used as a filtering medium, seems also well adapted for the purpose. If lime be precipitated by Clark's, or other process, it usually carries down the mud with it.

Rain water and the water of streams in mountain districts of hard igneous rock are generally nearly free from mineral constituents. This is the case with the Glasgow water from Loch Katrine, and the Thirlmere water which is to supply Manchester. Such water, if cold enough, and free from mud and organic impurity, is the best for almost every purpose in tanning. Most river water, however, and all spring water, is contaminated with mineral matter which it has dissolved out of the soil and rocks through which it has flowed. The principal of these mineral constituents are lime and magnesia. These occur both as sulphates and chlorides, and as hydric carbonates, or "bicarbonates." The sulphates and chlorides constitute "permanent" hardness, while that due to bicarbonates is called "temporary," from the fact that on boiling, half the carbonic acid is driven off, and the lime or magnesia is deposited as an insoluble neutral carbonate, thus softening the water. Any water which can be softened in this way by boiling may also be softened by the addition of a suitable quantity of lime, thus:—

This is Clark's process, and the chalk may either be separated by subsidence, which quickly takes place, or by a special filter (Porter-Clark). Thus the Bristol water, which from determinations by Mr. W. N. Evans, contains considerable temporary hardness and but little of permanent, may be almost completely softened by Clark's method. (For method of determining hardness and quantity of lime required, see p. 97).

The lime and magnesia constituting permanent hardness may be removed by the addition of sodic carbonate (soda ash or crystals); but this is expensive on a large scale, and as an equivalent quantity of sodic sulphate or chloride is left in the water, it is for most purposes of questionable advantage, though in some cases useful for the feed water of boilers. When employed for this purpose, the water should if possible be softened and settled before using, instead of adding the soda in the boiler itself, as is generally done. Soda is the active ingredient of many boiler compositions. For preventing furring, most tanning materials or even waste tan-liquors are very effective, and the danger of any corrosive action is lessened by the addition of a portion of soda ash. So far as is yet known, from the tanning point of view, it is hardly necessary to make any distinction between lime and magnesia, which may be considered simply as "hardness." A hard water probably softens dried hides more slowly, though it is possible that the observed difference may be due in many cases to the lower temperature of wells from which hard water is generally derived. In the actual limes, the hardness of the water can have no appreciable influence, though if sodium sulphide be used alone, a certain waste occurs from temporary hardness, which may render it advisable to add a little lime. It is in the washing of the hides from lime that the influence is first distinctly felt. If limey goods, after unhairing, are placed in a water with much temporary hardness, the same action occurs as in Clark's water softening process, and chalk is deposited in the surface of the hides, making them harsh and apt to "frize" or roughen the grain in scudding, and causing bad colour by combining with the organic acids of the liquors. The common, but not wholly satisfactory, expedient is to add a little lime, or better, a few pailfuls of lime liquor to the water before putting in the hides. The best plan is to use a properly softened water. Permanent hardness is not injurious in this way.

The hardness of water, and the dissolved carbonic acid which it contains, are, together with its temperature, the principal factors which determine whether a hide will plump or fall in it. Almost the only accurate investigation of this point has been made by W. Eitner ('Der Gerber,' iii. 183). He placed pieces of hide, unhaired by sweating, and quite flat and fallen, in water for 4 days at a temperature of 46° F. (8° C.), with the following results:—

The peculiarities which were shown by the hide pieces on removal from the water were maintained throughout the tanning, which was conducted in imitation of the German method, the hide being swollen and coloured through in weak birch-bark liquors, made with distilled water and acidified in each case with equal quantities of lactic acid, and finally laid away till tanned in a mixture of oak bark and valonia. No. 6, from magnesia sulphate, was the best; then No. 2; No. 3 was less good, but all the pieces from 1 to 6 were firm, close, and of good substance and texture, No. 1 having swelled well in the sour liquor. On the other hand, 7 and 8 scarcely swelled in liquor, but remained flat throughout, and were looser, thinner, and of finer fibre. From this experiment it is clear that while sulphates and carbonates exert a favourable influence on plumping, chlorides do the reverse, not only not plumping themselves, but placing the hides in an unfavourable condition for the plumping action of acids in the liquors. These experiments are quite borne out by the writer's experience in practice. The water at the Lowlights Tannery, which in dry weather is mostly obtained from beds of what was originally sea-sand, and which consequently contains a very abnormal proportion of chlorides (up to 68 pts. NaCl per 100,000), requires special and very careful management to make thick leather, notwithstanding its containing a considerable quantity of calcium and magnesium sulphates. These facts also indicate the importance of the thorough removal of salt from hides intended for sole-leather. Plumping is not a desirable thing in leather intended for dressing purposes, and it is possible that the use of a small percentage of salt in the liquors or wash waters might enable bating to be dispensed with. Like a bate, salt would dissolve a small proportion of hide substance (see p. 19). There is no practicable means of removing chlorides from water, but Eitner suggests the addition of a small quantity of sulphuric acid to water containing much temporary hardness (bicarbonates), by which it is converted into permanent (sulphates), which, as we have seen, plumps better. For this purpose about 2·8 oz. of ordinary English vitriol (sp. gr. 1·490) per 100 cub. ft. of water is required for each part of lime Ca(OH)₂ per 100,000 (see p. 97 for testing of water). A simpler guide is to add enough to purple, but not to redden litmus paper, even after moving it about in the water for some minutes. The acid must of course be well mixed by plunging. It must be borne in mind that Eitner's experiment was on sweated hides, and that with limed hide, which is kept plump by the dissolved lime retained in the hide, the conditions are different, and different results as regards carbonic acid and bicarbonates would probably be obtained. Both these would convert the lime in the hide into chalk, which is both insoluble and inert, and the hide would probably fall, at any rate till the lime was completely carbonated, while hides would remain plumpest in waters most free from substances capable of neutralising lime. One of the waters most effective in plumping limed hides is that of the river at Lincoln. Its hardness and contents in chlorine is, as compared with Lowlights water in dry weather,

Both waters have a considerable quantity of organic matter, and both owe their hardness in part to magnesia. From this we might conclude, what may be à priori expected, that the softer the water, the plumper limed hides remain in it. I am informed, however, by Mr. S. L. Evans, that in the Dartmoor water, which is very soft, but peaty, hides fall rapidly. In this case the colouring matters of the peat, which are of the nature of very weak acids, probably neutralise the lime. It may also be remarked, that wherever the conditions of putrefaction or decaying organic matter is present, hides rapidly fall, for the same reasons as they do in a bate.

While the injurious effect of bicarbonates on limed hide is matter of common experience, their influence on liquors and tanning is not so well understood. It is certain that they neutralise and combine with the organic acids of the liquors, and probably with some species of tannin, and as 1 part per 100,000 amounts to 1 oz. per 100 cub. ft., the acid required to neutralise a very hard water amounts to something considerable. It is well known that hard waters make bad tea, and the influence of hardness on the extraction of tannin is a subject well worthy of investigation, and which the writer hopes to examine.

On dyeing, at least as regards dye-woods, the influence of bicarbonates is distinctly favourable, and this is also stated to be true of woad, cochineal, and indigo-carmine.

Beside lime and magnesia salts, water may contain sulphates and chlorides of soda and potash; but not carbonates of these bases in presence of permanent hardness. In soft waters carbonates are sometimes present, and form carbonate of lime in limed hides. Hides are said to soften rapidly in such water. Alkaline sulphates are not known to have any injurious action, and chlorides have already been spoken of. Iron may be present in solution as bicarbonate, but not in any other form in presence of bicarbonate of lime. It is removed completely with the temporary hardness by Clark's process, or boiling. Iron is much more common merely in suspension, as mud, but is always objectionable. Most waters contain a little silicic acid and alumina, and some few considerable quantities. Such waters are said to harden leather, but the writer knows of no case where they are in use in England; and their occurrence is comparatively rare.

For comparison, analyses of a few spring and river waters are given on p. 89.

Analyses of various Waters.

CHAPTER VI.

METHODS OF CHEMICAL ANALYSIS FOR THE TANNERY.

It is assumed that the reader has an elementary knowledge of chemistry, and of the common manipulations of the laboratory; but at the risk of giving information which to many is already familiar, the principles that underlie those methods of testing which are most applicable to technical purposes must be briefly explained.

Standard Solutions.—If 40 grm. of pure caustic soda (NaHO) be dissolved in water, and a little tincture of litmus added, it will be coloured a bright blue. If hydrochloric acid be now added, drop by drop, the litmus will at last become purple, and a single drop more would turn it a bright red. At this point the liquid is neither acid nor alkaline, and if it be evaporated to dryness, nothing will be left but 58·5 grm. of common salt (NaCl), while 18 grm. of water will be formed and have escaped. We have therefore used exactly 36·5 grm. of pure HCl, and if we dissolve 40 grm. of caustic soda in 1 litre of water, and 36·5 grm. of pure HCl in another, equal parts of these liquids will always exactly neutralise each other, forming nothing but common salt and water. It will be obvious that if we have a soda solution of the strength named, we can find the amount of hydrochloric acid in any solution of unknown strength, by seeing how much of it is required to neutralise, say, 10 c.c. (= 0·4 grm. soda) of the known solution. Instead of 40 grm. of caustic soda, we may take 56 grm. of potash to the litre, and it will exactly neutralise an equal volume of the hydrochloric solution containing 36·5 grm. If, again, we make a solution containing 49 grm. of pure sulphuric acid (SO₄H₂) per litre, it will neutralise an exactly equal volume of either the soda or the potash solution, thus being precisely equivalent to the HCl solution. Such solutions are called normal, and any normal acid solution will neutralise an equal volume of any normal alkali, and vice versâ. For many purposes normal solutions are too strong, and solutions containing ¹/₁₀ of the quantities required for normal solution are preferable; such solutions are called decinormal. All solutions containing known quantities of chemicals, and intended for use in volumetric analysis, are called Standard solutions.

Indicators.—The tincture of litmus used to show when the solution is exactly neutral is called an indicator, and many materials are used in a similar way in different analytical processes. Thus the indigo solution in Löwenthal's process is an indicator. A more useful indicator than litmus for tannery purposes is Dr. Lunge's "methyl orange," which is indifferent to carbonic acid, and may therefore be used in the cold with solutions of alkaline carbonates; which are much more easily made and preserved than those of the caustic alkalies necessary with litmus. It is very sensitive to mineral acids, but not equally so to organic. It may be obtained of Messrs. Mawson and Swan, of Newcastle; and as a minute quantity only must be used for each test, it is really cheaper than litmus, and a few grm. will last a lifetime. It must be dissolved in water, and not more than 2 or 3 drops taken for each titration. (Titration signifies an estimation by means of a standard solution.) Other indicators will be named in connection with the analytical methods in which they are used.

Instruments.—To practically carry out analysis by standard solutions, measuring glasses are required. One or more flasks marked in the neck to hold exact quantities (Fig. 9), one at least, holding 1 litre, are indispensable. One or two graduated cylinders (Fig. 10), holding 100 c.c., and divided into tenths of c.c., are very useful, and it is well also to have one holding a litre, and provided with a stopper (Fig. 11). This is called a "test mixer," but is not absolutely essential.

Pipettes (Fig. 12) are tubes with a mark on the stem by which exact quantities of liquid can be taken. Several holding 5, 10, 20, and 25 c.c. are necessary, and one holding 10 c.c. and divided into tenths is advisable. Most important of all is the burette (Fig. 13). If only one is to be had, it must be a Mohr's burette with a glass tap, but as alkaline solutions are apt to set glass taps fast, it is well to have one with a tap, and another with a pinchcock (Fig. 14). They should hold 50 or 25 c.c. and be divided into tenths. The burette in use is fixed in a stand (Fig. 15) and filled up to the top of the graduation, and the quantity of solution delivered is then shown by the scale. It is usual to read by the under side of the hollow of the liquid, keeping the eye carefully level with it.

A chemical balance suitable for the preparation of standard solutions and general analytical use, is shown in Fig. 16. The beam is provided with steel or rock-crystal knife-edges at the centre, which are supported on agate planes, and similar edges a support the pans. Except at the moment of weighing, the beam, and in good balances the pans also (at b), are steadied by supports raised by turning the milled head c. The long pointer d moving over a scale, shows when the beam is horizontal; but the weighing is performed, not by waiting till the balance comes to rest, but by noting when the oscillations are equal on each side of the zero point. The weights, which should run from 50 grm. downwards, are usually of brass (preferably gilded) down to 1 grm., while the fractions to 0·01 grm. are of platinum foil. Milligrammes and fractions are weighed by a "rider" of wire weighing 0·01 grm., and moved along the beam (which is graduated for the purpose like a steelyard) by the arms e. A fair balance should turn distinctly with 0·001 grm., and a good one with 0·0001 grm. If equal weights are placed on each pan, they should of course balance, and if changed side for side the balance should be maintained. If not, the arms of the beam are unequal. Weights always have trifling errors, but if by a really good maker, these are generally so small that they may be disregarded except in very delicate researches. The weights should always be placed on the scale in regular order, beginning with the heaviest, and it is well to accustom oneself to reading the weight by the vacant places in the box as well as by the weights on the scale.

While of course it is most important, and for accurate work essential, to have as good a balance as possible, much may be done in technical work, even with a good pair of druggists' scales; and most standard solutions may be bought ready made; while from two or three accurately adjusted solutions many others may be made volumetrically.

Preparation of Standard Acid and Alkaline Solutions.—In practice it is very difficult to obtain perfectly pure caustic soda, free from water and carbonic acid, both of which are greedily absorbed by it from the air, so that a standard solution cannot practically be made by directly weighing out the substance as suggested in the introductory paragraph. In sodic carbonate, however, we have a substance which is easily obtained pure and dry, and which may be used for almost all the purposes to which a caustic solution could be applied. A decinormal solution is strong enough for most of the work in a tannery, though it is a convenience to have both normal and decinormal, and a stock of the stronger solution will last a longer time and is readily diluted to decinormal strength by adding 1 part to 9 parts of distilled water. To make a normal solution, about 60 grm. of the purest sodic carbonate are placed in a porcelain basin or platinum crucible and heated over a Bunsen gas-burner or spirit-lamp, nearly to redness, and allowed to cool closely covered up. Of the salt thus dried 53 grm. are accurately weighed into a beaker and dissolved in distilled water. The solution is then poured into a gauged litre flask, and carefully filled up with water at a temperature of 59° F. (15° C.) to the mark on the neck. The whole is then poured into a good-sized stoppered bottle (40 oz.) and vigorously shaken for 5-10 minutes. This thorough shaking is important with all standard solutions, and without experience no one would believe how much shaking is required uniformly to mix a solution. Probably more difficulty to beginners in analysis arises from neglect of this matter than from any other cause. To make a decinormal solution, proceed in precisely the same way, using 5·3 grm. instead of 53; or dilute as above.

Standard Acid Solution.—For this purpose any one of several acids may be used, each of which has its special advantages.

Oxalic acid is the easiest to make of any. A sufficient quantity of pure crystallised oxalic acid is powdered and pressed between filter paper, so as to absorb the moisture which occasionally is retained in cavities of the crystals. 6·3 grm. is then weighed out and dissolved in water, exactly as was done with sodic carbonate, forming a decinormal solution. It is used in Löwenthal's tannin estimation process and may also be employed to determine alkalies, but forms insoluble calcium oxalate with lime salts, and does not give a sharp reaction with methyl orange indicator. Hence litmus must be used, or a few drops of a neutral solution of calcium chloride added to the methyl orange, when hydrochloric acid will be liberated as soon as there is excess of the acid, and the indicator will be promptly reddened. Sulphuric acid is the most permanent of any acid solution, and may be generally employed. It forms insoluble sulphates with lime, baryta, and strontia. To make a normal solution, 35 c.c. of the pure concentrated acid are poured into at least 3 or 4 times as much distilled water, and allowed to cool, and are then made up to about 1 litre and well shaken. The burette is filled with the mixture, 10 c.c. of the standard sodic carbonate are measured into a beaker, 2 or 3 drops of methyl orange solution are added, and the acid is run in with constant stirring till the indicator is just beginning to redden. This must be repeated, and the two titrations should exactly agree. Suppose that 9.5 c.c. are required, then 950 c.c. of the trial acid are equal to 1 litre of the soda. If therefore 950 c.c. be measured into a test mixer, and made up to 1 litre, the solution should be accurately decinormal. Of course great care must be used in the whole process. If a gauged flask only is at hand it will be easier to measure into it the water required to make up the litre, and then fill to the mark with the trial acid. Normal hydrochloric acid may be made exactly as described for sulphuric acid, but using about 100 c.c. of the strongest acid. Decinormal solutions of both these acids may be made by the same methods; using 1 tenth the quantities, or by dilution of the normal solution.

Beside comparison with sodic carbonate solution, hydrochloric acid may also be checked by determining the amount of chlorine present, with silver nitrate (see p. 98) 10 c.c. of decinormal acid should of course be equal to 10 c.c. of decinormal silver nitrate.

Table giving the Quantity of the Following Substances contained in or equivalent to 1 litre of Normal or 10 litres of Decinormal Standard Solution.

EXAMINATION OF WATER.

Hardness (Hehner's process). (a) Temporary Hardness.—As has been stated (p. 84), this consists of lime and magnesia carbonates. As methyl orange is not affected by carbonic acid, bicarbonates of alkaline earths have an alkaline reaction, and may be estimated in solution by standard acid like the alkalies themselves. 100 c.c., or in soft waters 200 c.c., of the water is measured into a beaker, a drop or two of solution of methyl orange added, and decinormal hydrochloric or sulphuric acid run in from the burette with constant stirring till the colour just changes to pink. This is repeated, and the average taken. The two determinations should not at the most differ more than ¹/₁₀ c.c. Each c.c. represents 5 parts per 100,000 of CaCO₃ or 2·8 parts of CaO; or corresponding quantities of magnesia (4·2 parts of MgCO₃ or 2 parts MgO), when 100 c.c. of water are used.

(b) Permanent Hardness.—200 c.c. are measured into a beaker and boiled for 15 minutes with 40 c.c. decinormal sodic carbonate. The mixture is then allowed to cool and made up to 250 c.c.; or the flask and its contents may be weighed before boiling and made up again to the same weight. It is then filtered, and 60 c.c. representing 50 c.c. of the original water, is twice titrated with decinormal acid and the result added. If the water were pure, exactly 10 c.c. should be required to neutralise the 10 c.c. of sodic carbonate, but if there be permanent hardness a part of the sodic carbonate will be already neutralised with the acids of the lime and magnesia salts, which have been precipitated as carbonates together with the carbonates of these bases originally present in the water. The hardness will therefore be represented by the loss, i. e. the number of c.c. of acid used for 100 c.c. of the original water must be subtracted from 20 and the remainder calculated as before, or if calculated as sulphates, each c.c. represents 6·8 parts of CaSO₄ or 6 parts of MgSO₄ per 100,000. If, as is sometimes the case, more acid is required than is needed for the sodic carbonate used, the excess corresponds to sodic carbonate originally present in the water. In this case there can be no permanent hardness.

Chlorine in Water.—If silver nitrate be added to a solution of any chloride, the silver is precipitated as white curdy insoluble silver chloride. As indicator, a few drops of neutral potassic chromate are used. So long as any chloride is present the red silver chromate which forms is at once decomposed, and the silver converted into white chloride. But as soon as all the chloride is exhausted, the red chromate becomes permanent. To prepare a standard decinormal solution of silver, 17 grm. of pure recrystallised silver nitrate are dissolved in 1 litre of distilled water. To perform the estimation 50 c.c. of water are measured into a beaker, 2 or 3 drops of strong solution of pure yellow potassic chromate are added, and then silver nitrate from the burette till a permanent red is formed. This is repeated, and the results are added together, representing 100 c.c. of water. Each c.c. of silver nitrate used represents 3·55 parts of chlorine, or 5·85 parts of sodic chloride per 100,000. If more than 10 c.c. of silver solution are required to 50 c.c., it is advisable to use a smaller quantity of water. If the process be applied to other liquids than natural water, it must be borne in mind that the solution must not contain free acids or alkalies except carbonic acid. If this is not the case the liquid may be rendered faintly alkaline, with lime-water free from chlorides, and the excess of lime removed by passing carbonic acid through it; or it may be slightly acidified with sulphuric acid, and shaken with a little pure precipitated calcic or baric carbonate.

Detection of other Impurities.—Sulphuric acid (as sulphates) is seldom wholly absent, but its presence may be proved, by adding excess of barium chloride to the water slightly acidified with hydrochloric acid (2-3 c.c. of saturated solution of BaCl₂ are sufficient for any ordinary water); if the mixture be allowed to stand overnight in a 100 c.c. cylinder beside a solution containing a known, and not very different quantity of decinormal sulphuric acid, the quantity present may be roughly compared by measuring the bulk of the precipitates.

Lime may be similarly detected and roughly measured by precipitation with excess of ammonic oxalate in presence of ammonium chloride, to hinder precipitation of magnesia. Lime-water, which may be used as a standard, contains about 128 parts of lime per 100,000.

Magnesia is detected by adding ammonium phosphate to the filtrate from the precipitated oxalate of lime. If the mixture be allowed to stand in a warm place for 24 hours all the magnesia will be precipitated as ammonio-magnesic phosphate.

Silica, &c.—100 c.c. of the water is acidified with a little HCl evaporated to dryness, moistened with HCl, and treated with a little hot water. The silica or silicic acid is left undissolved. The solution from which the silicic acid has been filtered off is evaporated to small bulk and ammonia added, when iron will be precipitated as brown ferric oxide, which is coloured black by tannin or tanning liquor. If copper be present it will give a blue solution with the ammonia. Iron may also be recognised by evaporating the water to small bulk with a trace of HCl, and adding a little sodium acetate, when if iron be present it will be coloured black by tannin, red by ammonium sulphocyanide, and blue by potassium ferrocyanide (prussiate of potash). Its quantity may be estimated (Thomson, Chem. Soc. Abstracts, May 1885) by measuring 100 c.c. of the water to be tested and 100 c.c. distilled water into two similar cylinders, adding to each 5 c.c. of dilute hydrochloric acid (1:5) and 15 c.c. of a solution of potassium sulphocyanide (40 grm. per litre), and then adding to the distilled water cylinder a very dilute standard solution of ferric salt, till its colour matches the other. If the iron contained in the water is in the ferrous condition, it must be oxidised with potassic permanganate before testing.

A suitable ferric standard solution may be made by dissolving 0·1 grm. of clean, bright, soft iron wire in a little hydrochloric acid in a long-necked flask, adding nitric acid so long as red fumes are produced, evaporating nearly to dryness, and making up to 1 litre (more accurately 996 c.c.). Each c.c. will then equal 0·0001 grm. Fe.

Lead (and copper) may be detected by passing sulphuretted hydrogen through the water acidified with HCl, or by adding a drop of fresh ammonium or sodium sulphide to the slightly acidified water, when a brownish coloration clearly visible in a deep beaker set on a sheet of white paper will be produced. Iron also gives a black with sulphides in alkaline solution. Copper may be distinguished from lead by the blue given with ammonia, and by a reddish-brown precipitate with potassium ferrocyanide.

For accurate quantitative estimation of these impurities, the regular works on the subject, such as Thorpe's 'Quantitative Analysis,' Sutton's 'Volumetric Analysis,' or Fresenius' 'Quantitative Analysis,' must be consulted.

EXAMINATION OF COMMERCIAL ACIDS.

Sulphuric acid 10 grm. may be made up to 100 c.c. and well mixed, and of this 10 c.c. may be tested with normal sodic carbonate in presence of methyl orange. Each 1 c.c. of soda solution used corresponds to 0·049 grm. or 4·9 per cent. of H₂SO₄. For most purposes, the strength may be ascertained from the specific gravity, as measured by a hydrometer or weighed in a specific gravity bottle. The following table gives the strength at 59° F. (15° C.):—