 |
We may therefore use the following rules for working processes which do not yield proportional results. Make a series of two or three titrations, using very different quantities of metal in each. Subtract the lowest of these from the highest, and calculate the standard with the remainder. Calculate the volume required by this standard in any case, and find the excess or deficit, as the case may be. If an excess, subtract it from the result of each titration; if a deficit, add it; and use the standard in the usual way. The following table shows an actual example:—
- - Chalk taken. Gas obtained. Standard. - - 0.0873 gram 17.8 c.c. 0.4904 0.1305 " 27.3 " 0.4780 0.1690 " 35.8 " 0.4721 0.1905 " 40.4 " 0.4715 0.2460 " 52.5 " 0.4686 0.3000 " 64.0 " 0.4687 - -
It will be seen that the standard decreases as the quantity of chalk increases; this points to a deficiency in the quantity of gas evolved.
Then
0.3000 = 64.0 c.c. 0.0873 = 17.8 " ——— = —— 0.2127 = 46.2 "
and 0.2127100/46.2 = 0.4604. Then, multiplying the weight of chalk taken by 100, and dividing by 0.4604, we get the calculated results of the following table:—
+ + + -+ -+ Chalk taken. Gas found. Gas calculated. Difference. + + + -+ -+ 0.0873 gram 17.8 c.c. 18.9 c.c. -1.1 c.c. 0.1305 " 27.3 " 28.3 " -1.0 " 0.1690 " 35.8 " 36.7 " -0.9 " 0.1905 " 40.4 " 41.4 " -1.0 " 0.2460 " 52.5 " 53.4 " -0.9 " 0.3000 " 64.0 " 65.1 " -1.1 " + + + -+ -+
By adding 1 c.c. to the quantity of gas obtained, and taking 0.4604 as the standard, the calculated results will agree with those found with a variation of 0.1 c.c. When a large number of assays of the same kind are being made, this method of calculation is convenient; when, however, only one or two determinations are in question, it is easier to make a couple of standardisings, taking quantities as nearly as possible the same as those present in the assays.
Sometimes it is necessary to draw up a table which will show, without calculation, the weight of substance equivalent to a given volume of gas or of solution. The substance used for standardising should be, whenever possible, a pure sample of the substance to be determined—that is, for copper assays pure copper should be used, for iron assays pure iron, and so on; but when this cannot be got an impure substance may be used, provided it contains a known percentage of the metal, and that the impurities present are not such as will interfere with the accuracy of the assay. Including compounds with these, the standard may be calculated by multiplying the standard got in the usual way, by the percentage of metal in the compound or impure substance, and dividing by 100. If, for example, the standard 1.008 gram was obtained by using a sample of iron containing 99.7 per cent. of metal, the corrected standard would be 1.00899.7/100 = 1.005.
In volumetric analysis the change brought about must be one in which the end of the reaction is rendered prominent either by a change of colour or by the presence or absence of a precipitate. If the end of the reaction or finishing-point is not of itself visible, then it must be rendered visible by the use of a third reagent called an indicator.
For example, the action of sulphuric acid upon soda results in nothing which makes the action conspicuous; if, however, litmus or phenolphthalein be added the change from blue to red in the first case, or from red to colourless in the second, renders the finishing-point evident. Some indicators cannot be added to the assay solution without spoiling the result; in which case portions of the assay solution must be withdrawn from time to time and tested. This withdrawal of portions of the assay solution, if rashly done, must result in loss; if, however, the solution is not concentrated, and if the portions are only withdrawn towards the end of the titration, the loss is very trifling, and will not show-up on the result. The usual plan adopted is to have a solution of the indicator placed in drops at fairly equal intervals distributed over a clean and dry white porcelain-plate: a drop or two of the solution to be tested is then brought in contact with one of these and the effect noted. Another plan is to have thin blotting-paper, moistened with a solution of the indicator and dried; a drop of the solution to be tested placed on this shows the characteristic change. When the assay solution contains a suspended solid which interferes with the test, a prepared paper covered with an ordinary filter-paper answers very well; a drop of the solution to be tested is placed on the filter-paper, and, sinking through, shows its effect on the paper below.
Except when otherwise stated, all titrations should be made at the ordinary temperature; cooling, if necessary, by holding the flask under the tap. When a titration is directed to be made in a boiling solution, it must be remembered that the standard solution is cold, and that every addition lowers the temperature of the assay.
On running the solution from the burette into the assay, do not let it run down the side of the flask. If a portion of the assay has to be withdrawn for testing, shake the flask to ensure mixing, and then take out a drop with the test-rod; the neglect of these precautions may give a finishing-point too early. This is generally indicated by a sudden finish, in which case on shaking the flask and again testing no reaction is got. Do not remove the drop on the point of the burette with the test-rod; let it remain where it is or drop it into the solution by carefully opening the clip.
Generally the methods of working are as follows:—
(1) When the finishing-point depends on a change of colour in the solution.—Increase the bulk of the assay up to from 100 to 150 c.c. with water. Boil or cool, as the case may be. Run in the standard solution from a burette speedily, until the re-agent appears to have a slower action, and shake or stir all the time. Then run 1 c.c. or so at a time, still stirring, and finally add drops until the colour change is got.
(2) When an outside-indicator is used.—Pour the standard solution from a burette into the assay until 5 or 6 c.c. from the finishing-point; then run in 1 c.c. at a time (stirring and testing on the plate between each) until the indicator shows the change wanted, and deduct 0.5 c.c. for excess. When greater accuracy is sought for a duplicate assay is made. In this case the standard solution is run in close up to the end, and the operation is finished off with a few drops at a time.
(3) Where the finishing-point depends upon the absence of a precipitate and no outside-indicator is used.—As in the last case, run in the standard solution up to within a few c.c. of the end, then run in 1 c.c. at a time until a precipitate is no longer formed, but here 1.5 c.c. must be deducted for excess, since it is evident that the whole of the last "c.c." must have been, and a portion of the previous one may have been, in excess.
Indirect Titration.—The action of permanganate of potash upon a ferrous solution is one of oxidation, hence it is evident that if any other oxidising agent is present it will count as permanganate. In such a case the titration can be used (indirectly) to estimate the quantity of such oxidising agent, by determining how much less of the permanganate is used. For example, suppose that 1 gram of iron dissolved in sulphuric acid requires 100 c.c. of standard permanganate to fully oxidise it, but that the same amount of iron only requires 35.6 c.c. of the same standard permanganate if it has been previously heated with 0.5 gram of black oxide of manganese. Here it is evident that 0.5 gram of black oxide does the work of 64.4 c.c.[4] of the permanganate solution, and that these quantities are equivalent; moreover, if 64.4 c.c. correspond with 0.5 gram, then 100 c.c. correspond with 0.7764 which is the standard. On theoretical grounds, and by a method of calculation which will be explained further on (under the heading "Calculations from Formul"), it can be found that if the standard for iron is 1 gram, that for the black oxide will be 0.7764 gram.
The principles of these indirect titrations become clearer when expressed in a condensed form. Thus, in the example selected, and using the formul Fe = Iron, KMnO{4} = permanganate of potash, and MnO{2} = oxide of manganese, we have:—
(1) 1 gram Fe = 100 c.c. KMnO_{4}
(2) 1 gram Fe = 35.6 c.c. KMnO_{4} + 0.5 gram MnO_{2} .'. 100 c.c. KMnO_{4} = 35.6 c.c. KMnO_{4} + 0.5 gram MnO_{2} (100-35.6) c.c. KMnO_{4} = 0.5 gram MnO_{2} 64.4 c.c. KMnO_{4} = 0.5 gram MnO_{2}
The iron does not enter into the calculation if the same quantity is present in the two experiments.
An indirect titration thus requires three determinations, but if more than one assay is to be carried on, two of these need not be repeated. The standard is calculated in the usual way.
Colorimetric Assays.—These are assays in which the colour imparted to a solution by some compound of the metal to be determined is taken advantage of; the depth of colour depending on the quantity of metal present. They are generally used for the determination of such small quantities as are too minute to be weighed. The method of working is as follows:—A measured portion of the assay solution (generally 2/3, 1/2, 1/3, or 1/4 of the whole), coloured by the substance to be estimated, is placed in a white glass cylinder standing on a sheet of white paper or glazed porcelain. Into an exactly similar cylinder is placed the same amount of re-agents, &c., as the portion of the assay solution contains, and then water is added until the solutions are of nearly equal bulk. Next, a standard solution of the metal being estimated is run in from a burette, the mixture being stirred after each addition until the colour approaches that of the assay. The bulk of the two solutions is equalised by adding water. Then more standard solution is added until the tints are very nearly alike. Next, the amount added is read off from the burette, still more is poured in until the colour is slightly darker than that of the assay, and the burette read off again. The mean of the readings is taken, and gives the quantity of metal added. It equals the quantity of metal in the portion of the assay. If this portion was one-half of the whole, multiply by two; if one-third, multiply by three, and so on. When the quantity of metal in very dilute solutions is to be determined, it is sometimes necessary to concentrate the solutions by boiling them down before applying the re-agent which produces the coloured compound. Such concentration does not affect the calculations.
Gasometric Assays.—Gasometric methods are not much used by assayers, and, therefore, those students who wish to study them more fully than the limits of this work will permit, are recommended to consult Winkler and Lunge's text-book on the subject. The methods are without doubt capable of a more extended application. In measuring liquids, ordinary variations of temperature have but little effect, and variations of atmospheric pressure have none at all, whereas with gases it is different. Thus, 100 c.c. of an ordinary aqueous solution would, if heated from 10 C. to 20 C., expand to about 100.15 c.c. 100 c.c. of a gas similarly warmed would expand to about 103.5 c.c., and a fall of one inch in the barometer would have a very similar effect. And in measuring gases we have not only to take into account variations in volume due to changes in temperature and atmospheric pressure, but also that which is observed when a gas is measured wet and dry. Water gives off vapour at all temperatures, but the amount of vapour is larger as the temperature increases.
By ignoring these considerations, errors of 3 or 4 per cent. are easily made; but, fortunately, the corrections are simple, and it is easy to construct a piece of apparatus by means of which they may be reduced to a simple calculation by the rule of three.
The volume of a gas is, in practice, usually reduced to that which it would be at a temperature of 0 C., when the column of mercury in the barometer is 760 mm. high. But, although convenient, this practice is not always necessary. The only thing required is some way of checking the variations in volume, and of calculating what the corrected volume would be under certain fixed conditions.
Suppose that at the time a series of standardisings is being made, 100 c.c. of air were confined in a graduated tube over moist mercury. These 100 c.c. would vary in volume from day to day, but it would always be true of them that they would measure 100 c.c. under the same conditions as those under which the standardisings were made. If, then, in making an actual assay, 35.4 c.c. of gas were obtained, and the air in the tube measured 105 c.c., we should be justified in saying, that if the conditions had been those of the standardising, the 105 c.c. would have measured 100 c.c., and the 35.4 c.c. would have been 33.7; for 105: 100:: 35.4: 33.7. The rule for using such a piece of apparatus for correcting volumes is:—Multiply the c.c. of gas obtained by 100, and divide by the number of c.c. of air in the apparatus.
If it is desired to calculate the volumes under standard conditions (that is, the gas dry, at 0 C. and 760 mm. barometric pressure) the calculations are easily performed, but the temperature and pressure must be known.
Correction for Moisture.—The "vapour tension" of water has been accurately determined for various temperatures, and it may be looked upon as counteracting the barometric pressure. For example, at 15 C. the vapour tension equals 12.7 millimetres of mercury; if the barometer stood at 750 mm., the correction for moisture would be made by subtracting 12.7 from 750, and taking 737.3 mm. to be the true barometric pressure.
The vapour tensions for temperatures from 0 C. to 20 C. are as follows:—
-+ + -+ + -+ Temp. Tension. Temp. Tension. Temp. Tension. -+ + -+ + -+ 0 4.6 mm. 7 7.5 mm. 14 11.9 mm. 1 4.9 mm. 8 8.0 mm. 15 12.7 mm. 2 5.3 mm. 9 8.6 mm. 16 13.5 mm. 3 5.7 mm. 10 9.2 mm. 17 14.4 mm. 4 6.1 mm. 11 9.8 mm. 18 15.3 mm. 5 6.5 mm. 12 10.5 mm. 19 16.3 mm. 6 7.0 mm. 13 11.2 mm. 20 17.4 mm. -+ + -+ + -+ -
The correction for pressure is:—Multiply the volume by the actual pressure and divide by 760.
The correction for temperature:—Multiply the volume by 273 and divide by the temperature (in degrees Centigrade) added to 273.
For all three corrections the following rules hold good. To reduce to 0 C. and 760 mm. dry.
Volume 0.3592 (Pressure-tension) Corrected volume = ——————————————————— Temperature + 273
To find the volume, which a given volume under standard conditions would assume, if those conditions are altered.
Volume 2.784 (Temperature + 273) Resulting volume = —————————————————— Pressure - tension
As an example, we will suppose that it is desired to enclose in the apparatus referred to on p. 45, a volume of air, which, when dry (at 0 C. and 760 mm.), shall measure 100 c.c., whilst the actual temperature is 15 C., and the pressure 750 mm.
The second formula is the one to be used, and we get 108.7 c.c.
100 c.c.2.784288 Required volume = ——————————— 750-12.7
80179.2 = ———- 737.3
= 108.7 c.c.
FOOTNOTES:
[4] 100-35.6 = 64.4.
CHAPTER V.
WEIGHING AND MEASURING.
Weighing.—The system of weights and measures which we have adopted is the French or metric system; in this the gram (15.43 grains) is the unit of weight; the only other weight frequently referred to is the milligram, which is 0.001, or 1/1000 gram. The unit of volume is the cubic centimetre, which is approximately the volume of 1 gram of water, and which thus bears to the gram the same relation as grain-measures bear to grains. It is usual to write and even pronounce cubic centimetre shortly as c.c., and the only other denomination of volume we shall have occasion to use is the "litre," which measures 1000 c.c., and is roughly 1-3/4 pints.
The weights used are kept in boxes in a definite order, so that the weights on the balance can be counted as well by noting those which are absent from the box as by counting those present on the scale-pan. The weights run 50, 20, 10, 10, 5, 2, 1, 1 and 1 grams, and are formed of brass. The fractions of the gram are generally made of platinum or of aluminium, and are arranged in the following order:—0.5, 0.2, 0.1, 0.1, and 0.05, 0.02, 0.01, 0.01. These may be marked in this way, or they may be marked 500, 200, 100, 100, 50, 20, 10, 10; the 500 meaning 500 milligrams.
Some makers send out weights in the series 50, 20, 20, 10, &c.
Weights of less than 0.01 gram are generally present in a box, but it is much more convenient to work with a rider. This is a piece of wire which in the pan weighs 0.01 gram; it is made in such a form that it will ride on the beam, and its effective weight decreases as it approaches the centre. If the arm of the beam is divided into tenths, then each tenth counting from the centre outward equals 0.001 gram or 1 milligram, and if these tenths be further subdivided the fractions of a milligram are obtained; and these give figures in the fourth place of decimals. A fairly good balance should be sensitive to 0.0001 gram. The weights must never be touched with the fingers, and the forceps for moving them is used for no other purpose. When not in actual use the box is kept closed. The weights must not be allowed to remain on the pan of the balance. The balance-case must not be open without some reason. It must be fixed level, and, once fixed, must not be needlessly moved. The bench on which it stands should be used for no other purpose, and no one should be allowed to lean upon it.
When using a balance sit directly in front of it. Ordinarily the substance to be weighed is best put on the pan to the user's left; the weights and the rider are then easily manipulated. Powders, &c., should not be weighed directly on the balance; a counterpoised watch-glass or metal scoop (fig. 25) should be used. In some cases it is advisable to use a weighing-bottle. This is a light, well-stoppered bottle (fig. 3) containing the powdered ore. It is first filled and weighed; then some of the substance is carefully poured from it into a beaker or other vessel, and it is weighed again; the difference in the two weighings gives the weight of substance taken. A substance must always be cold when weighed, and large glass vessels should be allowed to stand in the balance-box a little while before being weighed. Always have the balance at rest when putting on or taking off anything from the pans. Put the weights on systematically. In using the rider (except you have a reason to the contrary), put it on at the 5; if this is too much, then try it at the 3; if then the weights are too little, try at the 4, if still not enough, the correct weight must be between the 4 and 5; try half-way between.
It is best to work with the balance vibrating; equilibrium is established when the vibration to the left is the mean of the preceding and succeeding vibrations to the right. For example, if it vibrates 6 divisions to the right on one swing, and 5 divisions on the next, the intermediate vibration to the left should have been 5-1/2.
Note whether the substance increases in weight whilst on the balance. If it does it may be because it was put on warm, and is cooling, or it may be because it is taking up moisture from the air. Substances which take up moisture rapidly should be weighed in clipped watch-glasses or in light-weighing bottles or tubes.
Students, in recording the weights, should first read off those missing from the box, writing down each order of figures as determined; first tens, then units, and so on. Remember that the first four platinum weights give the figures of the first place of decimals, the second four give the second place, and that the third and fourth places are given by the rider. Having taken down the figures, confirm them by reading off the weights as you put them back into the box. Do not rest a weight on the palm of your hand for convenience in reading the mark upon it. Remember one weight lost from a box spoils the set. Do not take it for granted that the balance is in equilibrium before you start weighing: try it.
Measuring Liquids.—For coarse work, such as measuring acids for dissolving ores, graduated glasses similar to those used by druggists may be used. It is well to have two sizes—a smaller graduated into divisions of 5 c.c. (fig. 26), and a larger with divisions equal to 10 c.c. No measurement of importance should be made in a vessel of this kind, as a slight variation in level causes a serious error.
Graduated flasks must be used when anything has to be made up to a definite bulk, or when a fixed volume has to be collected. If, for example, a certain weight of substance has to be dissolved and diluted to a litre, or if the first 50 c.c. of a distillate has to be collected, a flask should be used. Each flask is graduated for one particular quantity; the most useful sizes are 1000 c.c., 500 c.c., 200 c.c., 100 c.c., and 50 c.c. The mark should be in the narrowest part of the neck, and should be tangential to the curved surface of the liquid when the flask contains the exact volume specified. The level of a curved surface of liquid is at first somewhat difficult to read: the beginner is in doubt whether the surface should be taken at A, B, or C (fig. 27). It is best to take the lowest reading C. In some lights it is difficult to find this; in such cases a piece of white paper or card held behind and a little below, so as to throw light up and against the curved surface, will render it clear. In reading, one should look neither up at nor down upon the surface, but the eye should be on the same level with it. It must be kept in mind that flasks contain the quantity specified, but deliver less than this by the amount remaining in them and damping the sides. If it is desired to transfer the contents say of a 100 c.c. flask to a beaker, it will be necessary to complete the transfer by rinsing out the flask and adding the washings; otherwise there will be a sensible loss. Graduated cylinders (fig. 28) are convenient for preparing standard solutions.
Pipettes and burettes are graduated to deliver the quantities specified. The principle of the pipette, and the advantages and disadvantages of its various forms, may be understood by considering the first form shown in fig. 29. It is essentially a bulbed tube drawn out to a jet at its lower end, and having on each side of the bulb a mark so placed that when the surface of the liquid falls from the upper to the lower mark the instrument shall deliver exactly 100 c.c. The bore of the jet should be of such a size as will allow the level of the liquid to fall at the rate of about one foot in two minutes. If it runs more quickly than this, an appreciable error arises from the varying amount of liquid remaining, and damping the sides of the bulb. The flow of liquid from a pipette must not be hastened by blowing into it. The lower tube or nose of the pipette should be long enough to reach into the bottle or flask containing the liquid about to be measured. The pipette is filled by sucking at the open end with the mouth; this method of filling renders the use of the instrument dangerous for such liquids as strong acids, ammonia, and such poisonous solutions as that of potassic cyanide. One attempt with a fairly strong solution of ammonia will teach the beginner a very useful lesson. As soon as the liquid rises above the upper mark in the pipette, the mouth is withdrawn, and the pipette quickly closed by pressing the upper aperture with the index finger of the right hand; it is well to have the finger slightly moist, but not damp. The neck of the pipette should be long enough to allow its being firmly grasped by the fingers and thumb of the right hand without inconvenience. The pipette is first held in a vertical position long enough to allow any moisture outside the tube to run down, and then the liquid is allowed to run out to the level of the upper mark; this is easily effected by lessening the pressure. If the finger is wet, the flow will be jerky, and good work impossible. The pipette is next held over the vessel into which the 100 c.c. are to be put, and the liquid allowed to run out. When the bulb is nearly empty, the flow should be checked by replacing the finger, and the liquid allowed to escape slowly until the lower mark is reached. The pipette is then withdrawn; it is in the withdrawing that the disadvantage of this particular form[5] makes itself felt. It must be withdrawn very steadily, as the slightest shock causes the remaining column of liquid to vibrate, whereby air is drawn in and the liquid is forced out.
This disadvantage is got rid of by making the mouth of the jet the lower limit, or, in other words, allowing the instrument to empty itself. There are two forms of such pipettes; in the one generally recommended in Gay-Lussac's silver assay (the last shown in fig. 29) the nose is replaced by a jet. This is most conveniently filled by stopping the jet with the finger, and allowing the liquid to flow in a fine stream into the neck until the pipette is filled, and then working as just described. The other form is the one in general use; in fact, a long nose to a pipette is so convenient that it may almost be said to be necessary. But the accuracy is slightly diminished; a long narrow tube makes a poor measuring instrument because of the amount of liquid it finally retains. A defect possessed by both forms is the retention of a drop of varying size in the nozzle. Whatever method is adopted for removing this drop must be always adhered to. The most convenient form is the one last described, and the most useful sizes are 100 c.c., 50 c.c., 20 c.c., 10 c.c., and 5 c.c. Ten c.c. pipettes graduated into tenths of a cubic centimetre are very useful: those are best in which the graduation stops short of the bottom.
All measurements should be made at the ordinary temperature; and, before being used, the pipette should be rinsed out with a cubic centimetre or so of the solution to be measured. After using, it should be washed out with water.
Burettes differ mainly from pipettes in having the flow of liquid controlled from below instead of from above. The best form is that known as Mohr's, one kind of which is provided with a glass stopcock, while the other has a piece of india-rubber tube compressed by a clip. The latter cannot be used for solutions of permanganate of potash or of iodine, or of any substance which acts on india-rubber; but in other respects there is little to choose between the two kinds. A burette delivering 100 c.c., and graduated into fifths (i.e., each division = 0.2 c.c.), is a very convenient size. For some kinds of work, 50 c.c. divided into tenths (i.e., each division = 0.1 c.c.) may be selected.
Burettes may be fixed in any convenient stand; they must be vertical and should be so placed that the assayer can read any part of the graduated scale without straining. When not in use, they should be kept full of water. When using a burette, the water must be run out; the burette is next rinsed with some of the solution to be used, and drained; and then it is filled with the solution. Next squeeze the india-rubber tube so as to disentangle air-bubbles and, by smartly opening the clip, allow the tube and jet to be filled; see that no bubbles of air are left. Then run out cautiously until the level of the liquid in the burette stands at zero. In reading the level with very dark-coloured liquids it is convenient to read from the level A (fig. 27), and, provided it is done in each reading, there is no objection to this. The accuracy of the reading of a burette is sensibly increased by the use of an Erdmann float. This is an elongated bulb, weighted with mercury, and fitting (somewhat loosely) the tube of the burette. It floats in the solution, and is marked with a horizontal line; this line is taken as the level of the liquid. If the burette is filled from the top, the float rises with aggravating slowness, and this is its chief disadvantage. The float must come to rest before any reading is made.
A convenient plan for filling a burette from below is shown in fig. 30. The diagram explains itself. The bottle containing the standard solution is connected with the burette by a syphon arrangement through the glass tube and T-piece. The flow of liquid into the burette is controlled by the clip. When this clip is opened, the burette fills; and when it is closed, the burette is ready for use in the ordinary way.
Measuring Gases.—Lange's nitrometer (fig. 69) is a very convenient instrument for many gasometric methods. It requires the use of a fair quantity of mercury. In fig. 31, there is a representation of a piece of apparatus easily fitted up from the ordinary material of a laboratory. It is one which will serve some useful purposes. It consists of a wide-mouthed bottle fitted (by preference) with a rubber cork. The cork is perforated, and in the perforation is placed a glass tube which communicates with the burette. The burette is connected by a rubber tube and a Y-piece, either with another burette or with a piece of ordinary combustion-tube of about the same size. The wide-mouthed bottle contains either a short test-tube or an ordinary phial with its neck cut off. In working the apparatus the weighed substance is put in the bottle and the re-agent which is to act on it, in the test-tube; the cork is then inserted. The liquid in the two burettes is next brought to the same level, either by pouring it in at A or running it out at B. The level of the liquid in the apparatus for correcting variation in volume is then read and noted. Next, after seeing that the level of the liquid in the burette has not changed, turn the bottle over on its side so that the re-agent in the test-tube shall be upset into the bottle. Then, as the volume of the gas increases, lower the liquid in the burette by running it out at B, and at the same time keep the level in A half an inch or so lower than that in the burette. When the action has finished bring the liquid in the two vessels to the same level and read off the burette. This part of the work must always be done in the same manner.
The volume corrector for gas analysis is a graduated glass tube of 120 c.c. capacity inverted over a narrow glass cylinder of mercury. It contains 0.2 or 0.3 c.c. of water and a volume of air, which, if dry and under standard conditions, would measure 100 c.c. The actual volume varies from day to day, and is read off at any time by bringing the mercury inside and outside to the same level. This is done by raising or lowering the tube, as may be required. Any volume of gas obtained in an assay can be corrected to standard temperature and pressure by multiplying by 100 and dividing by the number of c.c. in the corrector at the time the assay is made.
FOOTNOTES:
[5] It is best to use this form with a glass stopcock, or with an india-rubber tube and clip, after the manner of a Mohr's burette.
CHAPTER VI.
RE-AGENTS.—ACIDS, ETC.
Acetic Acid, H[=A=c] or C_{2}H_{4}O_{2}. (sp. gr. 1.044, containing 33 per cent. real acid).—An organic acid, forming a class of salts, acetates, which are for the most part soluble in water, and which, on ignition, leave the oxide or carbonate of the metal. It is almost always used in those cases where mineral acids are objectionable. To convert, for example, a solution of a substance in hydrochloric acid into a solution of the same in acetic acid, alkali should be added in excess and then acetic acid. Many compounds are insoluble in acetic acid, which are soluble in mineral acids, such as ferric phosphate, ferric arsenate, zinc sulphide, calcium oxalate, &c., so that the use of acetic acid is valuable in some separations. The commercial acid is strong enough for most purposes, and is used without dilution.
"Aqua Regia" is a mixture of 1 part by measure of nitric acid and 3 parts of hydrochloric acid. The acids react forming what is practically a solution of chlorine.[6] The mixture is best made when wanted, and is chiefly used for the solution of gold and platinum and for "opening up" sulphides. When solutions in aqua regia are evaporated, chlorides are left.
Bromine, Br. (sp. gr. 3.0). Practically pure bromine.—It is a heavy reddish-brown liquid and very volatile. It boils at 60 C., and, consequently, must be kept in a cool place. It gives off brown irritating vapours, which render its use very objectionable. Generally it answers the same purpose as aqua regia, and is employed where the addition of nitric acid to a solution has to be specially avoided. It is also used for dissolving metals only from ores which contain metallic oxides not desired in the solution.
"Bromine Water" is simply bromine shaken up with water till no more is dissolved.
Carbonic Acid, CO_{2}.—A heavy gas, somewhat soluble in water; it is mainly used for providing an atmosphere in which substances may be dissolved, titrated, &c., without fear of oxidation. It is also used in titrating arsenic assays with "iodine" when a feeble acid is required to prevent the absorption of iodine by the alkaline carbonate. It is prepared when wanted in solution, by adding a gram or so of bicarbonate of soda and then as much acid as will decompose the bicarbonate mentioned. When a quantity of the gas is wanted, it is prepared, in an apparatus like that used for sulphuretted hydrogen, by acting on fragments of marble or limestone with dilute hydrochloric acid.
Citric Acid (H_{3}[=C=i] or C_{6}H_{8}O_{7}.H_{2}O) is an organic acid which occurs in colourless crystals, soluble in less than their weight of water. The solution must be freshly prepared, as it gets mouldy when kept. It forms a comparatively unimportant class of salts (citrates). It is used in the determination of phosphoric acid, chiefly for the purpose of preventing the precipitation of phosphates of iron and alumina by ammonia, and in a few similar cases. The commercial crystals are used; they should be free from sulphuric acid and leave no ash on ignition.
Hydrochloric Acid, HCl in water, (sp. gr. 1.16. It contains 32 per cent. of hydrogen chloride).—It is sometimes called "muriatic acid," and when impure, "spirit of salt." The acid solution should be colourless and free from arsenic, iron, and sulphuric acid. It forms an important family of salts, the chlorides. It is the best acid for dissolving metallic oxides and carbonates, and is always used by the assayer when oxidising agents are to be avoided. The acid is used without dilution when no directions are expressly given to dilute it. It has no action on the following metals: gold, platinum, arsenic, and mercury; it very slightly attacks antimony, bismuth, lead, silver, and copper. Tin is more soluble in it, but with difficulty; whilst iron, zinc, nickel, cobalt, cadmium, and aluminium easily dissolve with evolution of hydrogen and the formation of the lower chloride if the metal forms more than one class of salts. All the metallic oxides, except a few of the native and rarer oxides, are dissolved by it with the formation of chlorides of the metal and water.
Dilute Hydrochloric Acid is made by diluting the strong acid with an equal volume of water. This is used for dissolving precipitates obtained in the general course of analysis and the more easily soluble metals.
Hydrofluoric Acid, HF.—A solution in water may be purchased in gutta-percha or lead bottles. It is of variable strength and doubtful purity. It must always be examined quantitatively for the residue left on evaporation. It is used occasionally for the examination of silicates. It attacks silica, forming fluoride of silicon, which is a gas. When the introduction of another base will not interfere with the assay, the substance may be mixed in the platinum dish with fluoride of ammonium, or of potassium, or of calcium, and hydrochloric acid, instead of treating it with the commercial acid. It is only required in special work. The fumes and acid are dangerous, and, of course, glass or porcelain vessels cannot be used with it.
Iodine, I.—This can be obtained in commerce quite pure, and is often used for standardising. It is very slightly soluble in water, but readily dissolves in potassium iodide solution. It closely resembles chlorine and bromine in its properties, and can be used for dissolving metals without, at the same time, attacking any oxide which may be present. It is chiefly used as an oxidizing agent in volumetric work, being sharp in its reactions and easily detected in minute quantities. It cannot be used in alkaline solutions, since it reacts with the hydrates, and even with the carbonates, to form iodides and iodates. Iodine is soluble in alcohol.
Nitric Acid, HNO_{3}. (Sp. gr. 1.42; boiling point 121 C.; contains 70 per cent. by weight of hydrogen nitrate).—It is convenient to remember that one c.c. of this contains 1 gram of real acid. It combines the properties of an acid and of an oxidising agent. One c.c. contains 0.76 gram of oxygen, most of which is very loosely held, and easily given up to metals and other oxidisable substances. Consequently it will dissolve many metals, &c., upon which hydrochloric acid has no action. All sulphides (that of mercury excepted) are attacked by it, and for the most part rendered soluble. It has no action on gold or platinum, and very little on aluminium. The strong acid at the ordinary temperature does not act on iron or tin; and in most cases it acts better when diluted. Some nitrates being insoluble in nitric acid, form a protecting coat to the metal which hinders further action. Where the strong acid does act the action is very violent, so that generally it is better to use the dilute acid. When iron has been immersed in strong nitric acid it not only remains unacted on, but assumes a _passive_ state; so that if, after being wiped, it is then placed in the dilute acid, it will not dissolve. Tin and antimony are converted into insoluble oxides, while the other metals (with the exception of those already mentioned) dissolve as nitrates. During the solution of the metal red fumes are given off, which mainly consist of nitrogen peroxide. The solution is often coloured brown or green because of dissolved oxides of nitrogen, which must be got rid of by boiling. Generally some ammonium nitrate is formed, especially in the cases of zinc, iron, and tin, when these are acted on by cold dilute acid. Sulphur, phosphorus, and arsenic are converted into sulphuric, phosphoric, and arsenic acids respectively, when boiled with the strong acid.
Dilute Nitric Acid.—Dilute 1 volume of the strong acid with 2 of water.
Oxalic Acid, H_{2}Ō or (H_{2}C_{2}O_{4}.2H_{2}O.)—This is an organic acid in colourless crystals. It forms a family of salts—the oxalates. It is used in standardising; being a crystallised and permanent acid, it can be readily weighed. It is also used in separations, many of the oxalates being insoluble. For general use make a 10 per cent. solution. Use the commercially pure acid. On ignition the acid should leave no residue.
Sulphuretted Hydrogen. Hydrosulphuric acid, SH_{2}.—A gas largely used in assaying, since by its action it allows of the metals being conveniently classed into groups. It is soluble in water, this liquid dissolving at the ordinary temperature about three times its volume of the gas. The solution is only useful for testing. In separations, a current of the gas must always be used. It is best prepared in an apparatus like that shown in fig. 32, by acting on ferrous sulphide with dilute hydrochloric acid. When iron has to be subsequently determined in the assay solution, the gas should be washed by bubbling it through water in the smaller bottle; but for most purposes washing can be dispensed with. The gas is very objectionable, and operations with it must be carried out in a cupboard with a good draught. When the precipitation has been completed, the apparatus should always be washed out. The effect of this acid on solutions of the metals is to form sulphides. All the metallic sulphides are insoluble in water; but some are soluble in alkaline, and some in acid, solutions. If sulphuretted hydrogen is passed through an acid solution containing the metals till no further precipitation takes place, a precipitate will be formed containing sulphides insoluble in the acid. On filtering, adding ammonia (to render the filtrate alkaline), and again passing the gas, a further precipitate will be obtained, consisting of sulphides insoluble in an alkaline solution, but not precipitable in an acid one; the filtrate may also contain sulphides not precipitable in an acid solution, which are soluble in an alkaline one; these will be thrown down on neutralising. Again, the metals precipitated in the acid solution form sulphides which may be divided into groups, the one consisting of those which are soluble, and the other of those which are not soluble, in alkalies. This classification is shown in the following summary:—
1. Precipitable in an acid solution.
(a) Soluble in Alkalies.—Sulphides of As, Sb, Sn, Au, Pt, Ir, Mo, Te, and Se.
(b) Insoluble in Alkalies.—Sulphides of Ag, Pb, Hg, Bi, Cu, Cd, Pd, Rh, Os, and Ru.
2. Not precipitated in an acid solution, but thrown down in an alkaline one.
Sulphides of Mn, Zn, Fe, Ni, Co, In, Tl, and Ga.
These can again be divided into those which are dissolved by dilute acids and those which are not.
3. Not precipitated in an acid or alkaline solution, but thrown down on neutralising the latter.
Sulphides of V and W.
Sulphuretted hydrogen is a strong reducing agent. Ferric salts are thereby quickly reduced to ferrous; in hot solutions nitric acid is decomposed. These changes are marked by a precipitation of sulphur, and the student must be careful to pass the gas sufficiently long, and not be too hasty in concluding that no sulphide will form because it does not at once make its appearance. The best indication that it has been passed long enough is the smell of the gas in the solution after shaking.
Sulphurous Acid, H{2}SO{3}.—The reagent used may be regarded as a saturated solution of sulphur dioxide in water. It may be purchased, and keeps for a long time. It may be made by heating copper with sulphuric acid and passing the gas formed into water. The heat should be withdrawn when the gas is coming off freely. It is used as a reducing agent, and should not be diluted.
Sulphuric Acid, H{2}SO{4}. (Sp. gr. 1.84, containing 96 per cent. of real acid, H{2}SO{4}.)—This acid forms insoluble sulphates with salts of lead, strontium, and barium. It has a high boiling point, 290 C., and, when evaporated with salts of the more volatile acids, converts them into sulphates. When nitrates or chlorides are objectionable in a solution, evaporation with sulphuric acid removes them. In working with this acid caution is necessary, since, on mixing with water, great heat is evolved; and, if either the acid or water has been previously heated, a serious accident may result. In diluting the acid it should be poured into cold water. Glass vessels containing boiling sulphuric acid should be handled as little as possible, and should not be cooled under the tap. The action of diluted sulphuric acid on metals closely resembles that of dilute hydrochloric acid. Magnesium, aluminium, iron, zinc, nickel, cobalt, manganese, and cadmium dissolve, with evolution of hydrogen, in the cold acid, or when warmed. The action of hot and strong sulphuric acid is altogether different; it acts as an oxidising agent, and is itself reduced to sulphur dioxide or even to sulphur. The following metals are attacked in this way:—copper, bismuth, mercury, silver, antimony, tin, and lead. Gold, platinum, and arsenic are not affected. This property is made use of in parting silver from gold and platinum. Metallic sulphides are similarly attacked; but this method of opening up minerals has the disadvantage of giving rise to the formation of anhydrous sulphates of iron, &c., which are not readily dissolved when afterwards diluted. The use of sulphuric acid in assaying is (for these reasons) to be avoided. Its chief use is as a drying agent, since it has a strong affinity for water. Air under a bell jar may be kept dry by means of a basin of sulphuric acid, and gases bubbled through it are freed from water-vapour.
Dilute Sulphuric Acid.—This is made by diluting 1 volume of the strong acid with 4 of water.
Tartaric Acid, H{2}T or C{4}H{6}O{6}.—A crystallised organic acid, soluble in less than its own weight of water, or in less than three parts of alcohol. It is used for the same purposes as citric acid is. The solution is made when required.
BASES, SALTS, &c.
Alcohol, C{2}H{6}O. (Commercial alcohol of sp. gr. 0.838; it contains 84 per cent. by weight of alcohol.)—It should burn with a non-luminous flame and leave no residue. It is used for washing precipitates where water is inapplicable, and for facilitating drying.
Ammonia, NH_{3}. (Commercial ammonia, a solution having a sp. gr. of 0.88 to 0.89, and containing about 33 per cent. of ammonia.)—It is used as an alkali (more commonly than soda or potash), since an excess of it is easily removed by boiling. The salts of ammonium formed by it may be removed by igniting, or by evaporating in a porcelain dish with an excess of nitric acid. It differs in a marked way from soda or potash in its solvent action on the oxides or hydrates of the metals. Salts of the following metals are soluble in an ammoniacal solution in the presence of ammonic chloride:—copper, cadmium, silver, nickel, cobalt, manganese, zinc, magnesium, sodium, potassium, and the alkaline earths.
Dilute Ammonia is made by diluting 1 vol. of commercial ammonia with 2 of water. The dilute ammonia is always used; but in assays for copper a stronger solution (1 of strong ammonia to 1 of water) is required.
Ammonic Carbonate (Am{2}CO{3}) is prepared by dissolving one part of the commercial sesquicarbonate of ammonia in four parts of water, and adding one part of strong ammonia.
Ammonic Bicarbonate (HAmCO_{3}) is prepared by saturating a solution of the sesquicarbonate of ammonia with carbon dioxide.
Ammonic Chloride, AmCl.—Use the commercial salt in a 20 per cent. solution in water. The salt should leave no residue on ignition.
Ammonic Molybdate.—The solution is prepared as follows:—Dissolve 100 grams of the powdered commercial salt in 200 c.c. of dilute ammonia, and pour the solution in a slow stream into 750 c.c. of dilute nitric acid; make up to 1 litre, and allow the mixture to settle before using. It is used for the purpose of separating phosphoric oxide from bases and from other acids, and also as a test for phosphates and arsenates. In using this solution the substance must be dissolved in nitric acid, and a considerable excess of the reagent added (50 c.c. is sufficient to precipitate 0.1 gram P{2}O{5}); when the phosphate is in excess no precipitate will be got. The precipitate is phospho-molybdate of ammonia.
Ammonic Nitrate (AmNO_{3}) is used in the separation of phosphoric oxide by the molybdate method, and occasionally for destroying organic matter. It is soluble in less than its own weight of water. The solution is made when wanted.
Ammonic Oxalate (Am{2}C{2}O{4}.2H{2}O) is used chiefly for the separation of lime. The solution is made by dissolving 15 grams of the salt in 100 c.c. of water.
Ammonic Sulphide may be purchased in the state of a strong solution. It is yellow, and contains the disulphide, S{2}Am{2}. It serves the same purpose as is obtained by passing a current of sulphuretted hydrogen through an ammoniacal solution; but has the disadvantage of loading the solution with sulphur, which is precipitated when the solution is subsequently acidified. It is useful for dissolving the lower sulphide of tin (SnS).
Baric Carbonate (BaCO_{3}) is sometimes used for precipitating the weaker bases. It should be prepared when wanted by precipitating a solution of baric chloride with ammonic carbonate and washing. The moist precipitate is used without drying.
Baric Chloride, BaCl{2}.2H{2}O.—A crystallised salt, soluble in 2-1/2 parts of water. It is used for the detection and separation of sulphates. Make a 10 per cent. solution.
"Black Flux."—A mixture of finely divided carbon with carbonate of potash or with carbonates of potash and soda. It is prepared by heating tartar or "rochelle salt" until no more combustible gas is given off. One gram will reduce about 2 grams of lead from litharge.
Borax, Na{2}B{4}O{7}.10H{2}O.—It is chiefly used as a flux in dry assaying, as already described. It is also used in testing before the blowpipe; many metallic oxides impart a characteristic colour to a bead of borax in which they have been fused.
Calcium Chloride.—The crystallised salt is CaCl{2}.6H{2}O; dried at 200 C. it becomes CaCl{2}.2H{2}O, and when fused it becomes dehydrated. The fused salt, broken into small lumps, is used for drying gases. It combines with water, giving off much heat; and dissolves in a little more than its own weight of water. Strong solutions may be used in baths in which temperatures above the boiling-point of water are required. One part of the salt and 2 of water give a solution boiling at 112, and a solution of 2 parts of the salt in 1 of water boils at 158. The salt is very little used as a reagent.
Calcium Fluoride or "Fluor Spar," CaF_{2}.—The mineral is used as a flux in dry assaying; it renders slags which are thick from the presence of phosphates, &c., very fluid. Mixed with hydrochloric acid it may sometimes be used instead of hydrofluoric acid.
Calcium Carbonate, CaCO_{3}.—It is precipitated in a pure state by ammonic carbonate from a solution of calcium chloride. It is used for standardising. In the impure state, as marble or limestone, it is used in the preparation of carbonic acid.
Calcium Hydrate or "Lime Water."—This is used in testing for carbon dioxide and in estimating the amount of that gas present in air. It may be made by slaking quicklime and digesting the slaked lime with water. One hundred c.c. of water at 15 C. dissolves 0.1368 grams of the hydrate (CaH{2}O{2}), and hot water dissolves still less. "Milk of lime" is slaked lime suspended in water.
Cobalt Nitrate (Co(NO_{3})_{2}.6H_{2}O) is used in a 10 per cent. solution for the detection of oxides of zinc, aluminium, &c.; on ignition with which it forms characteristically coloured compounds.
Copper, Cu.—Pure copper, as obtained by electrolysis, can be purchased. This only should be used.
Copper Oxide, CuO.—It occurs as a black, heavy, and gritty power, and is used for the oxidation of carbon and hydrogen in organic substances. It should be ignited and cooled out of contact with air just before using, since it is hygroscopic. Oxide of copper which has been used may be again utilised after calcination.
Copper Sulphate (CuSO{4}.5H{2}O) contains 25.4 per cent. of copper. It is used in the outer cell of a Daniell-battery. The commercial salt is used for this purpose. The re-crystallised and pure salt is used for preparing the anhydrous sulphate, which is used for detecting moisture in gases. For this purpose it is dried at 200 C. till no trace of green or blue colour remains. It must be prepared when wanted. It may be conveniently used in the form of pumice-stone, saturated with a solution of the salt and dried. Traces of moisture develop a green colour.
Ferric Chloride, Fe_{2}Cl_{6}. (When crystallised, Fe_{2}Cl_{6}.6H_{2}O.)—The solution is prepared as described under iron. The commercial salt contains arsenic, and, since the chief use of ferric chloride is for the determination of this substance, it must be purified (_see_ under ARSENIC).
Ferric Sulphate (Fe{2}(SO{4}){3}) is a yellowish white deliquescent salt. It is used as an indicator in volumetric silver assaying, and for the separation of iodine from bromine. It may be purchased as iron alum, Am{2}Fe{2}(SO{4}){4}.24H{2}O. But it is best prepared by adding strong sulphuric acid to ferric hydrate in equivalent proportions. Use it as a solution containing 2 or 3 per cent. of iron.
Ferrous Sulphate, FeSO{4}.7H{2}O.—The granulated form is best, and can be purchased pure. It is used for standardising. It keeps better in crystals than in solution. It is readily soluble in water, but the solution is best made with the help of a little free acid. As a re-agent use a 10 per cent. solution. The crystals should be clear bluish-green; if their colour is dark green, brown, or blue, they should be rejected.
Ferrous Sulphide (FeS) is used for the preparation of sulphuretted hydrogen. It may be purchased and broken in small lumps, nut-size, for use.
"Fusion Mixture" (K{2}CO{3}.Na{2}CO{3}) is a mixture of potassic and sodic carbonates in the proportions of 13 of the former to 10 of the latter, by weight. It is hygroscopic. A mixture of the bicarbonates is better, being purer and less apt to get damp.
Gallic Acid (C{7}H{6}O{5}.H{2}O) is an organic acid, occurring as a pale fawn-coloured crystalline powder, soluble in 100 parts of cold water, or in 3 parts of boiling water. It is used for the determination of antimony. A 10 per cent. solution in warm water is made when required.
Hydrogen (H) is a gas. It is obtained by acting on zinc with dilute hydrochloric or sulphuric acid. It is used as a reducing agent, and for providing an atmosphere free from oxygen. It reduces metallic oxides at a high temperature. It must be freed from water; and special precautions should be taken to prevent an admixture with air. It is generally required in a current which can be continued for an hour or more without interruption. The preparation can be conveniently carried out in the apparatus shown (fig. 33). A quart bottle is half filled with sheet zinc, and connected with bulbs filled with sulphuric acid, and with a calcium chloride tube. The last is connected with the apparatus through which the gas has to be passed. Dilute hydrochloric acid mixed with a few cubic centimetres (20 c.c. to 1 pint) of stannous chloride sol. to fix any dissolved oxygen, is placed in the funnel, and let into the bottle by opening the stopcock when required. Care must be taken to let the hydrogen escape for some time before starting the reduction.
Gold, Au.—Gold, obtained by cupelling and "parting," is for most purposes sufficiently pure. It is best kept in the shape of foil. When the purer metal is required, gold should be dissolved in aqua regia, the solution evaporated to a paste, diluted, allowed to stand, and filtered. The filtered solution is acidified with hydrochloric acid, warmed, and precipitated with sodium sulphite. The precipitate is collected, washed, and fused on charcoal.
Iron, Fe.—The soft wire (thin) is used for standardising. Rods are used in dry assays as a desulphurising agent. Steel must not be used, since it is not pure, and contains a variable amount of iron.
Lead, Pb.—Granulated lead or lead-foil is used in the dry assay for silver and gold, and in the preparation of lead salts. It can be obtained very pure, but always contains more or less silver, 1 or 2 milligrams in 100 grams. The amount of silver it contains must be determined and recorded.
Lead Acetate (Pb[=A=c]_{2}.3H_{2}O, or Pb(C_{2}H_{3}O_{2})_{2}.3H_{2}O) is used as a test, specially for the detection and estimation of sulphuretted hydrogen. Prepare a 10 per cent. solution for use.
Lead Nitrate (Pb(NO{3}){2}) can be purchased pure. It is used for standardising.
Lead Dioxide (PbO_{2}) occurs as a dark-brown powder. It is used as an oxidizing agent and for absorbing sulphurous oxide. It can be prepared by digesting red lead with warm dilute nitric acid; washing and drying the residue.
"Litharge," PbO.—It can be purchased as a yellow heavy powder. It is used in dry assaying as a flux, as a desulphurising agent, and also as a source of lead. It always contains some silver, the amount of which must be determined.
Litmus.—This is an organic colouring matter which is turned red by acids and blue by alkalies. For ordinary purposes it is best used as litmus paper, which may be purchased in small books. A solution is prepared by digesting 15 or 20 grams of the commercial litmus in 100 c.c. of water on the water bath. After being allowed to settle, it is filtered and made just faintly red with acetic acid. Then there is added a drop or two of a solution of soda and 10 c.c. of alcohol. It should be kept in a loosely-covered bottle.
Magnesia, MgO.—It may be purchased as "calcined magnesia." It is used for making "magnesia mixture," and should be kept in a corked wide-mouthed bottle.
"Magnesia Mixture."—Dissolve 22 grams of magnesia in about a quarter of a litre of dilute hydrochloric acid, avoiding excess. Add 5 grams of magnesia, boil, and filter. Add 300 grams of ammonic chloride, and 250 c.c. of strong ammonia; and dilute with water to 2 litres. It should be kept in a stoppered winchester.
Magnesium Sulphate, MgSO{4}.7H{2}O.—It can be purchased very pure, and is occasionally used as a standard salt.
Manganese Dioxide, MnO_{2}.—It is used in the preparation of chlorine. The commercial article is not pure, but is sufficiently so for this purpose.
Marble, CaCO_{3}.—Fragments of the white crystalline variety only should be used. It is used as a source of lime and of carbon dioxide.
Mercury, Hg.—This can be purchased pure. It should have a bright surface, flow without a tail, and leave no residue on ignition. It is used as a standard; for amalgamation; and as a confining liquid in gas analysis.
Mercuric Chloride (HgCl_{2}) may be purchased pure. Make a 5 per cent. solution in water. It is used for destroying an excess of stannous chloride; for removing sulphuretted hydrogen from solution; and as a test for stannous salts.
Microcosmic Salt, HAmNaPO_{4}.8H_{2}O.—When fused NaPO_{3} is formed. It is used in testing for metallic oxides and silica before the blowpipe. The crystals are sometimes used as a standard for phosphoric acid.
"Nessler's Solution."—Mode of preparation: Dissolve 35 grams of potassium iodide in 100 c.c. of water; dissolve 17 grams of mercuric chloride in 300 c.c. of water, and pour this solution into that of the iodide till a permanent precipitate is produced; make up to 1 litre with a 20 per cent. solution of potash; add mercuric chloride till a precipitate is again formed; allow to settle and decant. It is used for detecting ammonia.
Nitre.—This is potassium nitrate.
Platinum Chloride, 2HCl.PtCl{4}. (In the crystallised form it has 6H{2}O).—It may be made as follows:—Take 5 grams of clean platinum scrap and dissolve in a flask at a gentle heat in 50 c.c. of hydrochloric acid with the occasional addition of some nitric acid; evaporate to a paste; and then dissolve in 100 c.c. of water. It is used for separating and determining potassium.
Phenolphthalein is an organic compound used as an indicator; more especially in determining the weaker acids, it cannot be used in the presence of ammonia. Dissolve half a gram in 100 c.c. of dilute alcohol.
Potassium Bicarbonate, KHCO_{3}.—It may be purchased pure; on ignition it leaves the carbonate, K_{2}CO_{3}, which may be used as a standard.
Potassium Cyanide, KCN.—It is used in the dry assay as a reducing agent. The commercial salt is very impure. Purchase that sold as potassic cyanide (gold) which contains about 95 per cent. of KCN. It is used for copper assaying and occasionally in separation. Make a 10 per cent. solution when wanted.
Potassium Bichromate, K_{2}Cr_{2}O_{7}. It may be purchased nearly pure. It is used as an oxidising agent, for determining iron; and as a test solution. For this last purpose a 10 per cent. solution is prepared.
Potassium Chlorate (KClO_{3}) can be purchased pure. It is used with hydrochloric acid as a substitute for aqua regia.
Potassium Ferrocyanide (K_{4}Fe(CN)_{6}.3H_{2}O), or "yellow prussiate of potash," is used as a test; as an indicator; and for the determination of zinc. Make a 5 per cent. solution.
Potassium Ferricyanide (K_{6}Fe_{2}(CN)_{12}), or "red prussiate of potash," is used for testing; and as an indicator. Make a 5 per cent. solution when wanted, as it decomposes on keeping.
Potassium Hydrate, KHO. Purchase that purified with alcohol. It is an alkali, and is used for absorbing carbonic acid, &c.
Potassium Iodide, KI. It may be purchased nearly pure. It is used as a test and for dissolving iodine. It should be used in a 10 per cent. solution freshly made. The solution decomposes on exposure to light, with separation of iodine.
Potassium Nitrate (KNO_{3}) can be purchased pure. It is used in the dry way as an oxidizing agent. It is very fusible. It decomposes at a low temperature into potassium nitrite (KNO_{2}) and free oxygen; and at a higher temperature leaves potash (K_{2}O). It oxidizes sulphur and carbon with explosive violence. This action may be moderated by mixing the nitre with carbonate of soda, common salt, or some other inert body.
Potassium Nitrite, KNO_{2}.—The commercial article is not pure, but is sufficiently so for the purpose required. A saturated solution is used in the separation of cobalt; the solution is made when wanted.
Potassium Permanganate, KMnO_{4}.—This salt can be purchased sufficiently pure. It is much used as an oxidizing agent.
Potassium Bisulphate (KHSO_{4}) is used as a dry reagent for opening up minerals. It fuses; and at a much higher temperature is converted into potassium sulphate with loss of sulphuric acid.
Potassium Sulphocyanate (KCNS) is used for the detection and determination of traces of ferric iron; as also in the separation of silver and copper from some of the other metals. Make a 10 per cent. solution. It should show no colour on the addition of hydrochloric acid.
"Red Lead" (Pb{3}O{4}) is used in the dry assay as a flux instead of litharge, from which it differs in containing a little more oxygen. When acted on by nitric acid a brown residue of lead dioxide is left, nitrate of lead going into solution. Like litharge it always carries silver; about 2 milligrams in 100 grams.
Silver, Ag.—Pure silver in foil is required as a standard. It may be prepared as follows:—Dissolve scrap silver in dilute nitric acid and decant off from any residue; dilute the solution with hot water and add hydrochloric acid until there is no further precipitate, stir; allow the precipitate to settle; decant and wash; dry the precipitate, mix it with twice its bulk of carbonate of soda and fuse the mixture in a crucible until tranquil; clean the button and roll or hammer it into foil.
Sodium Acetate, NaC{2}H{3}O{2}.3H{2}O.—The crystals may be purchased sufficiently pure. Make a 20 per cent. solution in water. It is used for replacing mineral acids by acetic acid.[7]
Sodium Acetate and Acetic Acid.—A solution is used in the determination of phosphates and arsenates; 100 grams of the salt is dissolved in 500 c.c. of acetic acid, and diluted with water to one litre.
Sodium Bicarbonate (NaHCO_{3})is used as a flux in dry methods. On ignition it leaves the carbonate (Na_{2}CO_{3}), which is used as a standard reagent. Make a 20 per cent. solution of the carbonate for use. It should be free from chlorides or sulphates, or if impure the amount of impurities must be determined.
Sodium Hydrate, NaHO. It may be purchased in sticks, which should be kept in a well-corked bottle. It is sometimes called "caustic soda." It is a strong alkali. It is used for neutralizing acid solutions and for separations where ammonia is unsuitable. Make a 5 per cent. solution for use.
Sodium Hyposulphite, Na{2}S{2}O{8}.5H{2}O.—It may be purchased pure. It is generally known as "hypo." It is used as a standard.
Sodium Sulphite (Na_{2}SO_{3}.7H_{2}O) is used as a reducing agent.
Sodium Phosphate, Na_{2}HPO_{4}.12H_{2}O. The crystals may be purchased pure, but they effloresce in dry air with loss of water. It is used as a standard and for precipitating magnesia, &c. Make a 10 per cent. solution.
Stannous Chloride, SnCl{2}.2H{2}O.—The crystals are best purchased. If kept dry and free from air they are fairly permanent. A solution is made by dissolving 20 grams in 10 c.c. of hydrochloric acid and diluting to 1 litre. The solution is not permanent. It is a strong reducing agent, and is chiefly used in solution for this purpose.
Tin, Sn.—Grain tin should be purchased. It is not pure, but contains 99.5 per cent. of the metal. The chief impurity is copper. It can be used as a standard. When acted on with hot hydrochloric acid it slowly dissolves (more rapidly in contact with platinum) and forms stannous chloride.
Uranium Acetate, UO{2}(C{2}H{3}O{2}){2}.H{2}O.—It is best purchased in crystals. The solution is used for the determination of phosphates and arsenates. A solution of 3 per cent. strength is occasionally used as an indicator.
Uranium Nitrate, UO{2}(NO{3}){2}.6H{2}O.—This salt is very soluble in water and is sometimes used instead of the acetate, which is somewhat difficult to dissolve.
"Water," H_{2}O.—Spring or well water is sufficiently pure for most purposes, 100 c.c. will leave a residue of from 10 to 30 milligrams, so that where a salt has to be dissolved out, evaporated, and weighed it should be replaced by distilled water. Rain water, melted snow, &c., always leave less residue than spring water; but in other respects they are often dirtier. Distilled water is best prepared in the office, a glass or tin condenser being used.
Zinc, Zn.—It is sold in a granulated form or in sticks. It generally contains over 1 per cent. of lead, with a little iron and arsenic. It is used for separating metals from their solutions, and generally as a reducing agent. For the preparation of hydrogen, and in most other cases, scrap sheet zinc may be used.
Zinc Oxide, ZnO.—The commercial oxide sometimes contains carbonate.
Zinc Sulphate, ZnSO{4}.7H{2}O.—It is occasionally used as a standard, and can be purchased nearly pure.
FOOTNOTES:
[6] 3HCl + HNO_{3} = Cl_{2} + NOCl + 2H_{2}O.
[7] NaC{2}H{3}O{2} + HCl = H{4}C{2}O{2} + NaCl.
CHAPTER VII.
FORMUL, EQUATIONS, ETC.
Formul and equations are a kind of short hand for expressing briefly and in the language of the atomic theory the facts of chemical composition and reaction. The convenience of this method of expressing the facts justifies a short description of it here.
On comparing the percentage composition of a series of compounds the proportions in which the elements combine appears to be regulated by no simple law. For example:
Realgar. Orpiment. Mispickel. Pyrites. Arsenic 71.4 60.9 46.0 — Sulphur 28.6 39.1 19.6 53.3 Iron — — 34.4 46.7 ——— ——— ——— ——— 100.0 100.0 100.0 100.0
But if in these examples the composition is calculated, not on 100 parts, but on 107, 246, 163, and 120 parts respectively, evidence of a simple law becomes apparent.
Realgar. Orpiment. Mispickel. Pyrites. Arsenic 75.0 150.0 75.0 — Sulphur 32.0 96.0 32.0 64.0 Iron — — 56.0 56.0 ——— ——— ——— ——— 107.0 246.0 163.0 120.0
It will be seen that the proportion of arsenic is 75 or twice 75, that of iron is 56, and that of sulphur 32 or some simple multiple of 32. The series of examples might be extended indefinitely, and it would still be found that the "combining proportions" held good. The number 75 is spoken of as the "combining weight," or, more frequently, as the "atomic weight" of arsenic. Similarly 56 is the atomic weight of iron, and 32 the atomic weight of sulphur. The importance of this law of chemical combination is altogether independent of the atomic theory; but this theory furnishes the simplest explanation of the facts. According to it a chemical compound is made up of exactly similar groups of particles. The particles of each elementary substance are all alike, but differ from those of other elements in weight. Ultimate particles are called atoms, and the groups of atoms are called molecules. The atomic weight of any particular element is the weight of its atom compared with the weight of an atom of hydrogen. The atom of sulphur, for instance, is 32 times as heavy as the atom of hydrogen, and the atomic weight of sulphur is 32. The molecular weight is the sum of the atomic weights of the group. The molecule of pyrites contains two atoms of sulphur and one of iron: on referring to the table of atomic weights it will be seen that the atomic weights are—sulphur 32, and iron 56. The molecular weight, therefore, is 32+32+56—that is, 120. The meaning of this is, 120 parts by weight of iron pyrites contain 64 parts of sulphur and 56 parts of iron; and this is true whether the "parts by weight" be grains or tons.
The symbol or formula of an atom is generally the initial letter or letters of the Latin or English name of the substance. The atom of hydrogen is written H, that of oxygen O, of sulphur S, of iron (ferrum) Fe, and so on. A list of these symbols is given in the table of atomic weights.
The formula of a molecule is obtained by placing together the symbols of the contained atoms. Thus, Fe represents an atom of iron, S an atom of sulphur, while FeS represents the molecule of sulphide of iron as containing one atom of each element.
When more than one atom of an element is present this is shown by writing a figure under and after the symbol; thus, FeS_{2} represents a molecule with one atom of iron and two atoms of sulphur, Fe_{2}S_{3} similarly shows one with two atoms of iron and three of sulphur. When a group of atoms is enclosed in brackets, a figure after and under the bracket multiplies all within it; for example, Pb(NO_{3})_{2} is another way of writing PbN_{2}O_{6}. Sometimes it is convenient to represent the atoms of a molecule as divided into two or more groups; this may be done by writing the formul of the groups, and separating each simple formula by a full stop. Slaked lime, for instance, has the formula CaH_{2}O_{2}; or, as already explained, we may write it Ca(HO)_{2}; or, if for purposes of explanation we wished to look on it as lime (CaO) and water (H_{2}O), we could write it CaO.H_{2}O. A plus sign (+) has a different meaning; CaO + H_{2}O indicates quantities of two substances, water and lime, which are separate from each other. The sign of equality (=) is generally used to separate a statement of the reagents used from another statement of the products of the reaction; it may be translated into the word "yields" or "becomes." The two statements form an equation.
Ignoring the quantitative relation, the meaning of the equation CaO + H{2}O = CaO.H{2}O is: "lime and water yield slaked lime." By referring to a table of atomic weights we can elicit the quantitative relations thus:—
CaO + H{2}O = CaH{2}O{2} V V V Ca = 40 H{2} = 2 = 12 Ca = 40 O = 16 O = 16 H{2} = 2 = 12 O{2} = 32 = 162 56 18 74
Or, putting it in words, 56 parts of lime combine with 18 parts of water to form 74 parts of slaked lime. This equation enables one to answer such a question as this:—How much lime must be used to produce 1 cwt. of slaked lime? for, if 74 lbs. of slaked lime require 56 lbs. of lime, 112 lbs. will require (56112)/74, or about 84-3/4 lbs.
As another example having a closer bearing on assaying take the following question:—"In order to assay 5 grams of 'black tin' (SnO_{2}) by the cyanide process, how much potassic cyanide (KCN) will be required?" The reaction is
SnO{2} + 2KCN = Sn + 2KCNO V V Sn = 118 K = 39 O{2} = 32 C = 12 - N = 14 150 652 = 130
What is sought for here is the relation between the quantities of SnO_{2} and KCN. Note that a figure before a formula multiplies all that follows up to the next stop or plus or equality sign. The question is now resolved to this: if 150 grams of oxide of tin require 130 grams of cyanide, how much will 5 grams require?
150 : 130 :: 5 : x x = 4.33 grams.
A problem of frequent occurrence is to find the percentage composition of a substance when its formula has been given. For example: "What percentage of iron is contained in a mineral having the formula 2Fe{2}O{3}.3H{2}O?" Bringing this formula together we have Fe{4}H{6}O{9}. Find the molecular weight.
Fe_{4} = 224 = 564 H_{6} = 6 = 16 O_{9} = 144 = 169 —- 374
Then we get: 374 parts of the mineral contain 224 of iron. How much will 100 contain?
374 : 224 :: 100 : x x = 59.89.
And the answer to the question is 59.89 per cent.
Again, suppose the question is of this kind:—"How much crystallised copper sulphate (CuSO{4}.5H{2}O) will be required to make 2 litres of a solution, 1 c.c. of which shall contain 0.0010 gram of copper?"
A litre is 1000 c.c., so, therefore, 2 litres of the solution must contain 0.001 gram 2000, or 2 grams. How much crystallised copper sulphate will contain this amount of metal?
Cu = 63.3 S = 32.0 O{4} = 64.0 = 164 5H{2}O = 90.0 = 185 ——- 249.3
If 63.3 grams of copper are contained in 249.3 grams of sulphate, in how much is 2 grams contained.
63.3 : 249.3 :: 2 grams : x x = 7.8769 grams.
The answer is, 7.8769 grams must be taken.
As a sample of another class of problem similar in nature to the last (but a little more complicated) take the following:—"What weight of permanganate of potash must be taken to make 2 litres of a solution, 100 c.c. of which shall be equivalent to 1 gram of iron?" In the first place the 2 litres must be equivalent to 20 grams of iron, for there are 20 100 c.c. in two litres. In the titration of iron by permanganate solution there are two reactions. First in dissolving the iron
Fe + H{2}SO{4} = FeSO{4} + H{2} V 56
and second, in the actual titration,
10FeSO_{4} + 2KMnO_{4} + 9H_{2}SO_{4}= 2MnSO_{4} + 5Fe_{2}(SO_{4})_{3} + 2KHSO_{4} + 8H_{2}O V K = 39 Mn = 55 O_{4}= 64 - 158 2 = 316
As before, attention is confined to the two substances under consideration—viz., Fe and KMnO_{4}. In the second equation, we find 316 parts of the permanganate are required for 10 molecules of FeSO_{4}; and in the first equation 56 parts of iron are equivalent to one molecule of FeSO_{4}, therefore 560 of iron are equivalent to 316 of permanganate; and the question is, How much of the permanganate will be equivalent to 20 grams of iron?
560 : 316 :: 20 grams : x. x= 11.286 grams.
The answer is 11.286 grams.
Very similar to this last problem is the question suggested under the head "Indirect Titration" (p. 43). "If 100 c.c. of the standard permanganate solution are equivalent to 1 gram of iron, how much peroxide of manganese will they be equivalent to?" The equation for dissolving the iron is already given; the second equation is
2FeSO{4} + MnO{2} + 2H{2}SO{4} = Fe{2}(SO{4}){2} + MnSO{4} + 2H{2}O V Mn = 55 O{2} = 32 87
It will be seen that 87 grams of peroxide of manganese are equivalent to 112 grams of iron. How much then is equivalent to 1 gram of iron?
112 : 87 :: 1 gram : x x = 0.7767 gram.
It is sometimes convenient to calculate the formula of a substance from its analysis. The method of calculating is shown by the following example. Required the formula of a mineral which gave the following figures on analysis:—
Cupric oxide (CuO) 10.58 Ferrous oxide (FeO) 15.69 Zinc oxide (ZnO) 0.35 Sulphuric oxide (SO{2}) 28.82 Water (H{2}O) 44.71 ——— 100.15
First find the molecular weights of CuO, FeO, &c., and divide the corresponding percentages by these figures. Thus, CuO = 63.3+16 = 79.3 and 10.58 divided by 79.3 gives 0.1334. Similarly FeO = 56+16 = 72 and 15.69 divided by 72 gives 0.2179. Treated in the same way the oxide of zinc, sulphuric oxide and water give as results 0.0043, 0.3602 and 2.484.
Classify the results as follows:—
Bases. Acids. Water.
CuO 0.1334 SO{3} 0.3602 H{2}O 2.484 FeO 0.2179 ZnO 0.0043 ————— ——————- —————— RO 0.3556 RO{3} 0.3602 R{2}O 2.484
The figures 0.3556, 0.3602 and 2.484 should be then divided by the lowest of them—i.e., 0.3556; or where, as in this case, two of the figures are very near each other the mean of these may be taken—i.e., 0.3579. Whichever is taken the figures got will be approximately 1, 1 and 7. The formula is then RO.SO{3}.7H{2}O in which R is nearly 2/5ths copper, 3/5ths iron and a little zinc.
This formula requires the following percentage composition, which for the sake of comparison is placed side by side with the actual results.
Calculated. Found. Cupric oxide 11.29 10.58 Ferrous oxide 15.37 15.69 Zinc oxide nil 0.35 Sulphuric oxide 28.47 28.82 Water 44.84 44.71 ——- ——— 99.97 100.15
Trimming the results of an analysis to make them fit in more closely with the calculations from the formula would be foolish as well as dishonest. There can be no doubt that the actual analytical results represent the composition of the specimen much more closely than the formula does; although perhaps other specimens of the same mineral would yield results which would group themselves better around the calculated results than around those of the first specimen analysed. It must be remembered that substances are rarely found pure either in nature or in the arts; so that in most cases the formula only gives an approximation to the truth. In the case of hydrated salts there is generally a difficulty in getting the salt with exactly the right proportion of water.
PRACTICAL EXERCISES.
The following calculations may be made:—
1. Calculate standards in the following cases— (a) Silver taken, 1.003 gram. Standard salt used, 100.15 c.c. (b) Iron taken, 0.7 gram. Bichromate used, 69.6 c.c.
2. Calculate percentages:— (a) Ore taken, 1 gram. Solution used, 65.2 c.c. Standard, 0.987 gram.
(b) Ore taken, 1 gram. Barium sulphate got, 1.432 gram. Barium sulphate contains 13.73 per cent. of sulphur, and the percentage of sulphur in the ore is wanted.
(c) Barium sulphate is BaSO_{4}. Calculate the percentage of sulphur it contains, for use in the preceding question.
3. A method of estimating the quantity of peroxide in a manganese ore is based on the following reactions:—
(1) MnO{2} + 4HCl = MnCl{2} + Cl{2} + 2H{2}O.
(2) Cl + KI = KCl + I.
To how much MnO_{2} is 1 gram of Iodine (I) equivalent?
4. A mineral has the following composition:—
Carbonic acid (CO{2}) 19.09 Copper oxide (CuO) 71.46 Water (H{2}O) 9.02
What is its formula?
5. How much copper is contained in 1.5 gram of crystallized copper sulphate (CuSO{4}.5H{2}O)? How much of these crystals must be taken to give 0.4 gram of copper?
6. How much ferrous sulphate crystals (FeSO{4}.7H{2}O) must be taken to yield 2 litres of a solution, 100 c.c. of which shall contain 0.56 gram of iron?
7. Galena is PbS, and hmatite Fe{2}O{3}. What percentages of metal do these minerals contain?
CHAPTER VIII.
SPECIFIC GRAVITY.
The relation of the weight of a substance to its volume should be kept in mind in all cases where both weight and volume are dealt with. Students are apt to imagine that on mixing equal volumes of, say, sulphuric acid and water, an acid of half the strength must be obtained. If the statement of strength is in parts by weight this will lead to considerable error. For example, 100 c.c. of sulphuric acid containing 98 per cent. by weight of real acid, will, if diluted with 100 c.c. of water, yield a solution containing not 49 per cent. by weight, but about 63.5 per cent. of the acid. The reason is this: the 100 c.c. of sulphuric acid weighs 184 grams, and contains 180.32 grams of real acid, while the 100 c.c. of water weighs only 100 grams; the mixed water and acid weighs 284 grams, and contains 180.32 of real acid, which is equivalent to nearly 63.5 per cent. by weight. If, however, the method of statement be volumetric, it would be correct to say that doubling the volume halves the strength: if 100 c.c. of brine contains 10 grams of salt, and is diluted with water to 200 c.c., it would be of one-half the former strength, that is, 100 c.c. of the solution would contain 5 grams of salt.
This confusion is avoided by always stating the strengths as so many grams or "c.c." in 100 c.c. of the liquid. But obviously it would be advantageous to be able to determine quickly the weight of any particular substance corresponding to 1 c.c. or some other given volume. Moreover, in descriptions of processes the strengths of acids and solutions are frequently defined neither by their gravimetric nor volumetric composition, but by a statement either of specific gravity or of the degrees registered by Twaddell's or Beaum's hydrometer. Thus, in the description of the process of gold parting, one writer gives: "The acid should be of 1.2 specific gravity"; and another says: "The acid must not be stronger than 32 Beaum."
These considerations justify an account of the subject in such a work as this. And on other grounds the determination of a specific gravity is one of the operations with which an assayer should be familiar.
The meaning of "specific gravity" is present in the mind of every one who uses the sentence "lead is heavier than water." This is meaningless except some such phrase as "bulk for bulk" be added. Make the sentence quantitative by saying: "bulk for bulk lead is 11.36 times heavier than water," and one has the exact meaning of: "the specific gravity of lead is 11.36." A table of the specific gravities of liquids and solids shows how many times heavier the substances are than water.
It is better, however, to look upon the specific gravity (written shortly, sp. g.) as the weight of a substance divided by its volume. In the metric system, 1 c.c. of water at 4 C. weighs with sufficient exactness 1 gram; consequently, the sp. g., which states how many times heavier than water the substance is, also expresses the weight in grams of one c.c. of it. So that if a 100 c.c. flask of nitric acid weighs, after the weight of the flask has been deducted, 120 grams, 1 c.c. of the acid weighs 1.2 gram, and the sp. g. is 1.2. The specific gravity, then, may be determined by dividing the weight of a substance in grams by its volume in c.c.; but it is more convenient in practice to determine it by dividing the weight of the substance by the weight of an equal volume of water. And since the volumes of all substances, water included, vary with the temperature, the temperature at which the sp. g. is determined should be recorded. Even then there is room for ambiguity to the extent that such a statement as the following, "the specific gravity of the substance at 50 C. is 0.9010," may mean when compared with water at 50 C. or 4 C., or even 15.5 C. For practical purposes it should mean the first of these, for in the actual experiments the water and the substance are compared at the same temperature, and it is well to give the statement of results without any superfluous calculation. In the metric system the standard temperature is 4 C., for it is at this point that 1 c.c. of water weighs exactly 1 gram. In England, the standard temperature is 60 F. (15.5 C.), which is supposed to be an average temperature of the balance-room. The convenience of the English standard, however, is merely apparent; it demands warming sometimes and sometimes cooling. For most purposes it is more convenient to select a temperature sufficiently high to avoid the necessity of cooling at any time. Warming to the required temperature gives very little trouble.
Determination of Specific Gravity.—There is a quick and easy method of determining the density or sp. g. of a liquid, based upon the fact that a floating body is buoyed up more by a heavy liquid than by a light one. The method is more remarkable for speed than accuracy, but still is sufficiently exact. The piece of apparatus used for the purpose is endowed with a variety of names—sp. g. spindle, hydrometer, areometer, salimeter, alcoholimeter, lactometer, and so on, according to the special liquid upon which it is intended to be used. It consists of a float with a sinker at one end and a graduated tube or rod at the other. It is made of metal or glass. Generally two are required, one for liquids ranging in sp. g. from 1.000 to 2.000, and another, which will indicate a sp. g. between 0.700 and 1.000. The range depends on the size of the instrument. For special work, in which variations within narrow limits are to be determined, more delicate instruments with a narrower range are made.
In using a hydrometer, the liquid to be tested is placed in a cylinder (fig. 34) tall enough to allow the instrument to float, and not too narrow. The temperature is taken, and the hydrometer is immersed in the fluid. The mark on the hydrometer stem, level with the surface of the liquid, is read off. With transparent liquids it is best to read the mark under and over the water surface and take the mean.
The graduation of hydrometers is not made to any uniform system. Those marked in degrees Baum or Twaddell, or according to specific gravity, are most commonly used. The degrees on Baum's hydrometer agree among themselves in being at equal distances along the stem; but they are proportional neither to the specific gravity, nor to the percentage of salt in the solution. They may be converted into an ordinary statement of specific gravity by the following formul:—
Sp. g. = 144.3/(144.3-degrees Baum.)
or putting the rule in words, subtract the degrees Baum from 144.3, and divide 144.3 with the number thus obtained. For example: 32 Baum equals a sp. g. of 1.285.
144.3/(144.3-32) = 144.3/(112.3) = 1.285
This rule is for liquids heavier than water; for the lighter liquids the rule is as follows:—
Sp. g. = 146/(136 + degrees Baum.)
or in words divide 146 by the number of degrees Baum added to 136. For example: ammonia of 30 Beaum has a sp. g. of 0.880 (nearly).
146/(136+30) = 146/166 = 0.8795
A simple series of calculations enables one to convert a Beaum hydrometer into one showing the actual sp. g. Graduation, according to sp. g. is the most convenient for general purposes. In these instruments the distances between the divisions become less as the densities increase.
Twaddell's hydrometer is graduated in this way: Each degree Twaddell is 0.005 in excess of unity. To convert into sp. g. multiply the degrees Twaddell by 0.005, and add 1. For example: 25 Twaddell equals a sp. g. of 1.125.
25.005 = 0.125; + 1.000 = 1.125.
There is a practice which ignores the decimal point and speaks of a sp. g. of 1125 instead of 1.125. In some cases it is convenient, and inasmuch as no substance has a real sp. g. of much over 20, it can lead to no confusion. The figures expressed in this way represent the weight of a litre in grams.
Some hydrometers are graduated so as to show at a glance the percentage composition of the liquid they are intended to be used with. Gay-Lussac designed one to show the alcoholic strength of mixtures of alcohol and water; the construction of others upon the same principle is easy and perhaps useful. But when the principle is applied to complex liquids and mixed solutions, it is misleading.
The various methods of graduation ought all to give place to one showing a simple statement of the sp. g.
The method of determining sp. g. with the hydrometer is obviously inapplicable to the case of solids, and in the case of liquids it should not be used where exact figures are required. There are several other methods which may be used, but on the whole those with the specific gravity bottle are most convenient.
The specific gravity bottle (fig. 35) is a light flask of about 25 c.c. capacity, provided with a well-fitting perforated stopper. It is essentially a graduated flask, which measures a constant volume, but it does not much matter what the volume is.
In taking the sp. g. of a liquid (or, what is the same thing, a fused solid) there is wanted the weights (1) of the flaskful of water and (2) of the flaskful of the liquid. Dividing the second by the first gives the required sp. g. The actual weighings required are—
(1) of the dry and empty flask,
(2) of the flask filled with water, and
(3) of the flask filled with the liquid.
The weighing of the flask once made need not be often repeated. It is well to do so now and then for safety's sake; but one weighing will serve for a large number of determinations. The same remarks apply to the weighing of the bottle filled with water. The bottle is dried by rinsing out first with alcohol and afterwards with ether; ether is very volatile, and a short exposure in a warm place will soon drive off the little remaining about the sides. The ether vapour should be sucked out through a glass tube. See that the bore of the stopper is dry as well as the bottle. Let the dry bottle stand in the box of the balance for a minute or two before weighing. The weight is, strictly speaking, not that of the empty bottle, but of the bottle filled with air. The empty bottle would weigh from 20 to 30 milligrams less. Correcting for this would, in most cases, only make a difference in the fourth place of decimals,[8] so that it is better to ignore the error.
The weight of the flask filled with water is got by filling it with distilled water, and inserting the stopper. The excess of water will overflow at the margin and through the bore. The bottle is wiped with a soft, dry cloth, taking care not to squeeze or warm the bottle. The bottle will remain filled to the top of the stopper. It is allowed to stand in the balance box for a minute or two, and then weighed.
Distilled water, as stated, should be used; the use of ordinary water may increase the weight by 5 or 6 milligrams. Many waters, if they have not previously been boiled, give off bubbles of air which render the weighing worthless.
The temperature of the water is of greater importance; lowering the temperature 2 will increase the weight by 10 or 12 milligrams. A beaker of water may be warmed or cooled to the required temperature; then the bottle is filled from it, and quickly weighed. If the balance-room is cooler than the water, the latter will draw back into the bottle, and a few small bubbles of air will enter; but even in extreme cases this will only increase the weight by a very small fraction of a milligram. There is more trouble caused when the room is warmer, for the liquid then expands and protrudes as a drop resting on the top of the stopper. There will in this case be loss by evaporation, which in the case of the more volatile liquids, such as alcohol, is serious. To prevent this loss, as well as any that may arise by overflow, the stopper should be dilated above into a small cup, A (fig. 36), which may itself be stoppered. In a bottle of this kind the neck of the stopper is graduated, and the bottle is considered full when the liquid stands at the level of the mark in the neck. On inserting the stopper, the liquid rises into the cup, and is reduced to the level of the mark by absorption with pieces of filter-paper.
For most purposes, however, there is no need for cooling and allowing room for subsequent expansion. The assayer, as a rule, can select his own standard temperature, and may choose one which will always necessitate warming. It will be handier in this case to have a bottle with a thermometer stopper. Of the two types shown in fig. 37, that with the external thermometer tube (A) is more generally useful.
The bottle is filled at a lower temperature, and is then gently warmed so as to slowly raise the temperature to the required degree. The superfluous liquid is then at once wiped off, and the bottle cooled and weighed.
The weight of the flask filled with the liquid whose sp. g. has to be determined is ascertained in a similar way. Of course the temperature must be the same. If the liquid does not mix with water, the bottle should be dried before filling, but otherwise the flask need only be rinsed out two or three times with the liquid.
Having obtained the three weighings, deduct the weight of the bottle from each of the others to get the weights of the water and liquid respectively. Divide the latter by the former, the result shows the sp. g. As an example, take the following, in which a rather large sp. g. bottle was used:—
1. Weight of bottle 39.299 gram
2. Weight of bottle and water 81.884 "
3. Weight of bottle and paraffin 73.146 "
By subtracting 1 from 2 and 3 the result is as follows:—
81.884 grams 73.146 grams 39.299 " 39.299 " ——— ——— 42.585 of water. 33.847 of paraffin.
Divide the weight of the paraffin by that of the water—
42.585)33.8470(0.7948 29.8095 ———- .......
The sp. g. of the paraffin is 0.7948.
The sp. g. of a fusible solid may be obtained in the same way at a temperature some degrees above its fusing point.
The sp. g. of a solid in powder or gravel sufficiently fine to pass through the neck of the bottle is easily determined. If the bottle filled with water weighs 50 grams, and there is placed on the pan alongside of it 20 grams of a sand, the weight of the two together will of course be 70 grams. But if the sand is put in the bottle, it evidently displaces its own bulk of water; and if, on again weighing, the weight is found to be 62 instead of 70 grams, it is because the 20 grams of sand has displaced 8 grams of water. Bulk for bulk, the sand is 2-1/2 times as heavy.
In practice, the weight of the bottle filled with water will probably be already known; if not, it must be determined. A certain quantity, say 20 grams, of the powdered substance is then transferred carefully to the bottle. The bottle need not be dry inside, but its neck and outside must be. In making this transference a careful worker will make no loss, and the mode of working saves a little time. But it is better to weigh the dry flask; put into it 10 to 20 grams of the powder, and weigh again. The increase in weight gives accurately the weight of powder in the bottle. About two-thirds fill the bottle with distilled water, and mix with the powder by gentle shaking. Air bubbles will disentangle themselves, and rise to the surface of the water. Wash back anything adhering to the stopper with a jet of water, and fill the bottle almost to overflowing. Allow it to stand for a minute or so; replace the stopper; warm to the required temperature; take off the superfluous moisture; wipe and weigh. As an example, take the following:— |
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