LECTURE IV WATER: ITS CHEMISTRY AND PROPERTIES; IMPURITIES AND THEIR ACTION; TESTS OF PURITY-- Continued
In the last lecture, under the head of "Solution," I mentioned that some salts, some chemical substances, are more soluble in water than others, and that their solubilities under different circumstances of temperature vary in different ways. However, some salts and compounds are practically insoluble in water under any circumstances. We now arrive at the important result known to chemists as the precipitation of insoluble compounds from solutions. In order to define this result, however, we must, of course, first consider the circumstances of causation of the result. Let us take a simple case of chemical decomposition resulting in the deposition or precipitation of a substance from solution in the insoluble state. We will take a salt you are probably acquainted with—sulphate of copper, or bluestone, and dissolve it in water, and we have then the sulphate of copper in solution in water. Now suppose it is our desire to obtain from that solution all the copper by depositing it in some insoluble form. We may accomplish this in several different ways, relying on certain methods of decomposing that sulphate of copper. One of the simplest and most economical is that adopted in a certain so-called wet method of extracting copper. It is based on the fact that metallic iron has a greater tendency to combine in water solutions, with the acids of copper salts, than the copper has in those salts. We simply need to place some scraps of iron in the copper sulphate solution to induce a change which may be represented as follows: Copper sulphate, consisting of a combination of copper oxide with sulphuric acid, yields with iron, iron sulphate, a combination of iron oxide with sulphuric acid, and metallic copper. The metallic copper produced separates in the form of a red coating on the iron scraps. But we may also, relying on the fact that oxide of copper is insoluble in water, arrange for the deposition of the copper in that form. This we can do by adding caustic soda to a hot solution of copper sulphate, when we get the following change: Copper sulphate, consisting of a combination of copper oxide with sulphuric acid, yields with caustic soda, sulphate of soda, a combination of soda with sulphuric acid and oxide of copper. Oxide of copper is black, and so in this decomposition what is called a "black precipitate" of that oxide is produced on adding the caustic soda. But it might not suit us thus to deposit the copper from our solution; we might desire to remove the sulphuric acid from the copper sulphate, and leave the copper dissolved, say in the form of a chloride. We select, then, a compound which is a chloride, and a chloride of a metal which forms an insoluble combination with sulphuric acid—chloride of barium, say. On adding this chloride of barium to sulphate of copper solution, we get then a change which we might represent thus: Copper sulphate, consisting of a combination of copper oxide with sulphuric acid, yields with barium chloride, which is a combination of barium and chlorine, insoluble barium sulphate, a combination of barium oxide with sulphuric acid, and soluble copper chloride, a combination of copper and chlorine. This is called a double interchange. Now these are a few illustrations to show you what is meant by chemical decompositions. One practical lesson, of course, we may draw is this: We must have a care in dissolving bluestone or copper sulphate, not to attempt it in iron pans, and not to store or put verdigris into iron vessels, or the iron will be acted upon, and to some extent the copper salt will become contaminated with iron. It will now be clear to you that, as a solvent for bodies usually soluble in water, water that is perfectly pure will be most suitable and not likely to cause any deposition or precipitation through chemical decompositions, for there are no salts or other compounds in pure water to cause such changes. Such pure water is called soft water. But the term is only a comparative one, and water that is not quite, but nearly pure—pure enough for most practical purposes—is also called soft water. Now rain is the purest form of natural water, for it is a kind of distilled water. Water rises in vapour from the ocean as from a still, and the salt and other dissolved matters remain behind. Meeting cold currents of air, the vapours condense in rain, and fall upon the earth. After coming in contact with the earth, the subsequent condition of that water entirely depends upon the character, as regards solubility or insolubility, of the substances composing the strata or layers of earth upon which it falls, and through which it sinks. If it meets with insoluble rocks—for all rocks are not insoluble—it remains, of course, pure and soft, and in proportion as the constituents of rock and soil are soluble, in that proportion does the water become hard. We all know how dangerous acid is in water, causing that water to act on many substances, the iron of iron vessels, the lime in soil or rock, etc., bringing iron and lime respectively into solution. Now the atmosphere contains carbonic acid, and carbonic acid occurs in the earth, being evolved by decomposing vegetation, etc. Carbonic acid is also soluble to a certain, though not large extent, in water. As we shall see, water charged with carbonic acid attacks certain substances insoluble in pure water, and brings them into solution, and thus the water soon becomes hard. About the close of the last lecture, I said that lime is, to a certain extent, soluble in cold water. The solution is called lime-water; it might be called a solution of caustic lime. When carbonic acid gas first comes in contact with such a solution, chalk or carbonate of lime, which is insoluble in water, is formed, and the lime is thus precipitated as carbonate. Supposing, however, we continued to pass carbonic acid gas into that water, rendered milky with chalk powder, very soon the liquid would clear, and we should get once more a solution of lime, but not caustic lime as it was at first, simply now a solution of carbonate of lime in carbonic acid, or a solution of bicarbonate of lime. I will take some lime-water, and I will blow through; my breath contains carbonic acid, and you will see the clear liquid become milky owing to separation of insoluble carbonate of lime, or chalk. I now continue blowing, and at length that chalk dissolves with the excess of carbonic acid, forming bicarbonate of lime. This experiment explains how it is that water percolating through or running over limestone strata dissolves out the insoluble chalk. Such water, hard from dissolved carbonate of lime, can be softened by merely boiling the water, for the excess of carbonic acid is then expelled, and the chalk is precipitated again. This would be too costly for the softening of large quantities of water, the boiling process consuming too much coal, and so another process is adopted. Quicklime, or milk of lime, is added to the water in the proper quantity. This lime unites with the excess of carbonic acid holding chalk in solution, and forms with it insoluble chalk, and so all deposits together as chalk. By this liming process, also, the iron of the water dissolved likewise in ferruginous streams, etc., by carbonic acid, would be precipitated. To show this deposition I will now add some clear lime-water to the solution I made of chalk with the carbonic acid of my breath, and a precipitate is at once formed, all the lime and carbonic acid together depositing as insoluble chalk. Hence clear lime-water forms a good test for the presence of bicarbonates of lime or iron in a water. But water may be hard from the presence of other salts, other lime salts. For example, certain parts of the earth contain a great deal of gypsum, or natural sulphate of lime, and this is soluble to some extent in water. Water thus hardened is not affected by boiling, or the addition of lime, and is therefore termed permanently hard water, the water hardened with dissolved chalk being termed temporarily hard water. I have said nothing of solid or undissolved impurities in water, which are said to be in suspension, for the separation of these is a merely mechanical matter of settling, or filtration and settling combined. As a general rule, the water of rivers contains the most suspended and vegetable matter and the least amount of dissolved constituents, whereas spring and well waters contain the most dissolved matters and the least suspended. Serious damage may be done to the dyer by either of these classes of impurities, and I may tell you that the dissolved calcareous and magnesian impurities are the most frequent in occurrence and the most injurious. I told you that on boiling, the excess of carbonic acid holding chalk or carbonate of lime in solution as bicarbonate, is decomposed and carbonate of lime precipitated. You can at once imagine, then, what takes place in your steam boilers when such water is used, and how incrustations are formed. Let us now inquire as to the precise nature of the waste and injury caused by hard and impure waters. Let us also take, as an example, those most commonly occurring injurious constituents, the magnesian and calcareous impurities. Hard water only produces a lather with soap when that soap has effected the softening of the water, and not till then. In that process the soap is entirely wasted, and the fatty acids in it form, with the lime and magnesia, insoluble compounds called lime and magnesia soaps, which are sticky, greasy, adhesive bodies, that precipitate and fix some colouring matters like a mordant. We have in such cases, then, a kind of double mischief—(i) waste of soap, (ii) injury to colours and dyes on the fabrics. But this is not all, for colours are precipitated as lakes, and mordants also are precipitated, and thus wasted, in much the same sense as the soaps are. Now by taking a soap solution, formed by dissolving a known weight of soap in a known volume of water, and adding this gradually to hard water until a permanent lather is just produced, we can directly determine the consumption of soap by such a water, and ascertain the hardness. Such a method is called Clark's process of determination or testing, or Clark's soap test. We hear a great deal just now of soaps that will wash well in hard water, and do wonders under any conditions; but mark this fact, none of them will begin to perform effective duty until such hard water has been rendered soft at the expense of the soap. Soaps made of some oils, such as cocoa-nut oil, for example, are more soluble in water than when made of tallow, etc., and so they more quickly soften a hard water and yield lather, but they are wasted, as far as consumption is concerned, to just the same extent as any other soaps. They do not, however, waste so much time and trouble in effecting the end in view, and, as you know, "Time is money" in these days of work and competition. After making a soap test as described above, and knowing the quantity of water used, it is, of course, easy to calculate the annual loss of soap caused by the hardness of the water. The monthly consumption of soap in London is 1,000,000 kilograms (about 1000 tons), and it is estimated that the hardness of the Thames water means the use of 230,000 kilograms (nearly 230 tons) more soap per month than would be necessary if soft water were used. Of course the soap manufacturers around London would not state that fact on their advertising placards, but rather dwell on the victorious onslaught their particular brand will make on the dirt in articles to be washed, in the teeth of circumstances that would be hopeless for any other brand of soap! I have referred to the sticky and adhesive character of the compounds called lime soaps, formed in hard waters. Now in washing and scouring wool and other fibres, these sticky lime soaps adhere so pertinaciously that the fibres, be they of wool, silk, or any other article, remain in part untouched, impermeable to mordant or colouring matter, and hence irregular development of colour must be the consequence. Also an unnatural lustre or peculiar bloom may in parts arise, ruining the appearance of the goods. In some cases the lime soaps act like mordants, attracting colouring matter unequally, and producing patchy effects. In the dye-baths in which catechu and tannin are used, there is a waste of these matters, for insoluble compounds are formed with the lime, and the catechu and tannin are, to a certain extent, precipitated and lost. Some colours are best developed in an acid bath, such as Cochineal Scarlet, but the presence of the bicarbonate of lime tends to cause neutralisation of the acidity, and so the dyeing is either retarded or prevented. Such mordants as "red liquor" and "iron liquor," which are acetates of alumina and iron respectively, are also wasted, a portion of them being precipitated by the lime, thus weakening the mordant baths.
Ferruginous Impurities in Water.—Iron in solution in water is very objectionable in dyeing operations. When the iron is present as bicarbonate, it acts on soap solutions like the analogous lime and magnesia compounds, producing even worse results. In wool scouring, cotton bleaching, and other processes requiring the use of alkaline carbonates, ferric oxide is precipitated on the fibre. A yellowish tinge is communicated to bleached fabrics, and to dye bright and light colours is rendered almost out of the question. You may always suspect iron to be present in water flowing from or obtained directly out of old coal pits, iron mines, or from places abounding in iron and aluminous shales. Moreover, you sometimes, or rather generally, find that surface water draining off moorland districts, and passing over ochre beds, contains iron, and on its way deposits on the beds of the streamlets conveying it, and on the stones, red or brown oxide of iron. All water of this kind ought to be avoided in dyeing and similar operations. The iron in water from old coal pits and shale deposits is usually present as sulphate due to the oxidation of pyrites, a sulphuret or sulphide of iron. Water from heaths and moorlands is often acid from certain vegetable acids termed "peaty acids." This acidity places the water in the condition of a direct solvent for iron, and that dissolved iron may cause great injury. If such water cannot be dispensed with, the best way is to carefully neutralise it with carbonate of soda; the iron is then precipitated as carbonate of iron, and can be removed.
Contamination of Water by Factories.—You may have neighbours higher up the stream than yourselves, and these firms may cast forth as waste products substances which will cause immense waste and loss. Amongst these waste products the worst are those coming from chemical works, paper works, bleach works, etc. If the paper works be those working up wood pulp, the pollutions of effluent water will be about as noxious as they well can be. You will have gums and resins from the wood, calcium chloride from the bleach vats, acids from the "sours"; resin, and resin-soaps; there may also be alumina salts present. Now alumina, lime, resin, and resin-soaps, etc., precipitate dyestuffs, and also soap; if the water is alkaline, some of the mordants used may be precipitated and wasted, and very considerable damage done.
Permanent hardness in water, due to the presence of gypsum or sulphate of lime in solution, may be remedied by addition of caustic soda. Of course, if an alkaline water is objectionable in any process, the alkali would have to be neutralised by the addition of some acid. For use in boilers, water might thus be treated, but it would become costly if large quantities required such treatment. Water rendered impure by contaminations from dyehouses and some chemical works can be best purified, and most cheaply, by simple liming, followed by a settling process. If space is limited and much water is required, instead of the settling reservoirs, filtering beds of coke, sand, etc., may be used. The lime used neutralises acids in the contaminated and impure water, precipitates colouring matters, mordants, soap, albuminous matters, etc.
Tests of Purity.—I will now describe a few tests that may be of value to you in deciding as to what substances are contaminating any impure waters that may be at hand.
Iron.—If to a water you suspect to be hard from presence of carbonate of lime or carbonate of iron in solution in carbonic acid, i.e. as bicarbonates, you add some clear lime-water, and a white precipitate is produced, you have a proof of carbonate of lime—hardness. If the precipitate is brownish, you may have, also, carbonate of iron. I will now mention a very delicate test for iron. Such a test would be useful in confirmation. If a very dilute solution of such iron water be treated with a drop or two of pure hydrochloric acid, and a drop or so of permanganate of potash solution or of Condy's fluid, and after that a few drops of yellow prussiate of potash solution be added, then a blue colour (Prussian blue), either at once or after standing a few hours, proves the presence of iron.
Copper.—Sometimes, as in the neighbourhood of copper mines or of some copper pyrites deposits, a water may be contaminated with small quantities of copper. The yellow prussiate once more forms a good test, but to ensure the absence of free mineral acids, it is first well to add a little acetate of soda solution. A drop or two of the prussiate solution then gives a brown colour, even if but traces of copper are present.
Magnesia.—Suppose lime and magnesia are present. You may first evaporate to a small bulk, adding a drop of hydrochloric acid if the liquid becomes muddy. Then add ammonia and ammonium oxalate, when lime alone is precipitated as the oxalate of lime. Filter through blotting paper, and to the clear filtrate add some phosphate of soda solution. A second precipitation proves the presence of magnesia.
Sulphates.—A solution of barium chloride and dilute hydrochloric acid gives a white turbidity.
Chlorides.—A solution of silver nitrate and nitric acid gives a white curdy precipitate.
Test for Lead in Drinking Water.—I will, lastly, give you a test that will be useful in your own homes to detect minute quantities of lead in water running through lead pipes. Place a large quantity of the water in a glass on a piece of white paper, and add a solution of sulphuretted hydrogen and let stand for some time. A brown colour denotes lead. Of course copper would also yield a brown coloration, but I am supposing that the circumstances preclude the presence of copper.
I have already said that rain water is the purest of natural waters; it is so soft, and free from dissolved mineral matters because it is a distilled water. In distilling water to purify it, we must be very careful what material we use for condensing the steam in, since it is a fact probably not sufficiently well known, that the softer and purer a water is, the more liable it is to attack lead pipes. Hence a coil of lead pipe to serve as condensing worm would be inadmissible. Such water as Manchester water, and Glasgow water from Loch Katrine still more so, are more liable to attack lead pipes than the hard London waters. To illustrate this fact, we will distil some water and condense in a leaden worm, then, on testing the water with our reagent, the sulphuretted hydrogen water, a brown colour is produced, showing the presence of lead. On condensing in a block tin worm, however, no tin is dissolved, so tin is safer and better as the material for such a purpose than lead.
Filtration.—We hear a great deal about filtration or filters as universal means of purifying water. Filtration, we must remember, will, as a rule, only remove solid or suspended impurities in water. For example, if we take some ivory black or bone black, and mix it with water and afterwards filter the black liquid through blotting-paper, the bone black remains on the paper, and clear, pure water comes through. Filtering is effective here. If we take some indigo solution, however, and pour it on to the filter, the liquid runs through as blue as it was when poured upon the filter. Filtering is ineffective here, and is so generally with liquids containing matters dissolved in them. But I said "generally," and so the question is suggested—Will filtration of any kind remove matters in solution? This question I will, in conclusion, try to answer. Bone charcoal, or bone black, has a wonderful attraction for many organic matters such as colours, dyes, and coloured impurities like those in peat water, raw sugar solutions, etc. For example, if we place on a paper filter some bone black, and filter through it some indigo solution, after first warming the latter with some more of the bone black, the liquid comes through clear, all the indigo being absorbed in some peculiar way, difficult to explain, by the bone black, and remaining on the filter. This power of charcoal also extends to gases, and to certain noxious dissolved organic impurities, but it is never safe to rely too much on such filters, since the charcoal can at length become charged with impurities, and gradually cease to act. These filters need cleaning and renewing from time to time.
