Fig. 33.
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]2.3H2O, or Pb(C2H3O2)2.3H2O) 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(NO3)2) can be purchased pure. It is used for standardising.
Lead Dioxide(PbO2) 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 withacetic 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, MgSO4.7H2O.—It can be purchased very pure, and is occasionally used as a standard salt.
Manganese Dioxide, MnO2.—It is used in the preparation of chlorine. The commercial article is not pure, but is sufficiently so for this purpose.
Marble, CaCO3.—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(HgCl2) 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, HAmNaPO4.8H2O.—When fused NaPO3is 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.PtCl4. (In the crystallised form it has 6H2O).—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.
Phenolphthaleinis 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, KHCO3.—It may be purchased pure; on ignition it leaves the carbonate, K2CO3, 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, K2Cr2O7. 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(KClO3) can be purchased pure. It is used with hydrochloric acid as a substitute for aqua regia.
Potassium Ferrocyanide(K4Fe(CN)6.3H2O), 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(K6Fe2(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(KNO3) 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 (KNO2) and free oxygen; and at a higher temperature leaves potash (K2O). 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, KNO2.—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, KMnO4.—This salt can be purchased sufficiently pure. It is much used as an oxidizing agent.
Potassium Bisulphate(KHSO4) is used as a dry reagent for opening up minerals. It fuses; and at a much higher temperatureis 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" (Pb3O4) 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, NaC2H3O2.3H2O.—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(NaHCO3)is used as a flux in dry methods. On ignition it leaves the carbonate (Na2CO3), 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, Na2S2O8.5H2O.—It may be purchased pure. It is generally known as "hypo." It is used as a standard.
Sodium Sulphite(Na2SO3.7H2O) is used as a reducing agent.
Sodium Phosphate, Na2HPO4.12H2O. 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, SnCl2.2H2O.—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, UO2(C2H3O2)2.H2O.—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, UO2(NO3)2.6H2O.—This salt is very soluble in water and is sometimes used instead of the acetate, which is somewhat difficult to dissolve.
"Water," H2O.—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, ZnSO4.7H2O.—It is occasionally used as a standard, and can be purchased nearly pure.
FOOTNOTES:[6]3HCl + HNO3= Cl2+ NOCl + 2H2O.[7]NaC2H3O2+ HCl = H4C2O2+ NaCl.
[6]3HCl + HNO3= Cl2+ NOCl + 2H2O.
[6]3HCl + HNO3= Cl2+ NOCl + 2H2O.
[7]NaC2H3O2+ HCl = H4C2O2+ NaCl.
[7]NaC2H3O2+ HCl = H4C2O2+ NaCl.
Formulæ 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.Arsenic71.460.946.0—Sulphur28.639.119.653.3Iron——34.446.7————————————100.0100.0100.0100.0
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.Arsenic75.0150.075.0—Sulphur32.096.032.064.0Iron——56.056.0————————————107.0246.0163.0120.0
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. Theparticles of each elementary substance are all alike, but differ from those of other elements in weight. Ultimate particles are calledatoms, and the groups of atoms are calledmolecules. 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. Themolecular weightis 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 atomis 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 moleculeis 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, FeS2represents a molecule with one atom of iron and two atoms of sulphur, Fe2S3similarly 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(NO3)2is another way of writing PbN2O6. 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 formulæ of the groups, and separating each simple formula by a full stop. Slaked lime, for instance, has the formula CaH2O2; 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 (H2O), we could write it CaO.H2O. A plus sign (+) has a different meaning; CaO + H2O 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 equationCaO + H2O = CaO.H2O is: "lime and water yield slaked lime." By referring to a table of atomic weights we can elicit the quantitative relations thus:—
CaO+H2O=CaH2O2↓↓↓Ca = 40H2= 2= 1×2Ca = 40O = 16O = 16H2= 2= 1×2————O2= 32= 16×25618——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 (56 × 112)/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' (SnO2) by the cyanide process, how much potassic cyanide (KCN) will be required?" The reaction is
SnO2+2KCN= Sn + 2KCNO↓↓Sn = 118K = 39O2= 32C = 12——N = 14150——65×2 = 130
What is sought for here is the relation between the quantities of SnO2and 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 :xx= 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 2Fe2O3.3H2O?" Bringing this formula together we have Fe4H6O9. Find the molecular weight.
Fe4= 224= 56×4H6= 6= 1×6O9= 144= 16×9——374
Then we get: 374 parts of the mineral contain 224 of iron. How much will 100 contain?
374 : 224 :: 100 :xx= 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 (CuSO4.5H2O) 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= 63.3S= 32.0O4= 64.0= 16×45H2O= 90.0= 18×5————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 :xx= 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 + H2SO4= FeSO4+ H2↓56
and second, in the actual titration,
10FeSO4+ 2KMnO4+ 9H2SO4= 2MnSO4+ 5Fe2(SO4)3+ 2KHSO4+ 8H2O↓K = 39Mn = 55O4= 64——158 × 2 = 316
As before, attention is confined to the two substances underconsideration—viz., Fe and KMnO4. In the second equation, we find 316 parts of the permanganate are required for 10 molecules of FeSO4; and in the first equation 56 parts of iron are equivalent to one molecule of FeSO4, 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
2FeSO4+ MnO2+ 2H2SO4= Fe2(SO4)2+ MnSO4+ 2H2O↓Mn = 55O2= 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 :xx= 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)10.58Ferrous oxide (FeO)15.69Zinc oxide (ZnO)0.35Sulphuric oxide (SO2)28.82Water (H2O)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 0.1334SO30.3602H2O 2.484FeO 0.2179ZnO 0.0043——————————————————————————————RO 0.3556RO30.3602R2O 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.SO3.7H2O 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 oxide11.2910.58Ferrous oxide15.3715.69Zinc oxidenil0.35Sulphuric oxide28.4728.82Water44.8444.71——————99.97100.15
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.
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 BaSO4. 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) MnO2+ 4HCl = MnCl2+ Cl2+ 2H2O.(2) Cl + KI = KCl + I.To how much MnO2is 1 gram of Iodine (I) equivalent?4. A mineral has the following composition:—Carbonic acid (CO2) 19.09Copper oxide (CuO) 71.46Water (H2O) 9.02What is its formula?5. How much copper is contained in 1.5 gram of crystallized copper sulphate (CuSO4.5H2O)? How much of these crystals must be taken to give 0.4 gram of copper?6. How much ferrous sulphate crystals (FeSO4.7H2O) 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 hæmatite Fe2O3. What percentages of metal do these minerals contain?
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 Beaumé'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° Beaumé."
These considerations justify an account of the subject in such a work as this. And on other grounds the determination of a specificgravity 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 dividingthe 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 forspeed 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.
Fig. 34.
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 Baumé or Twaddell, or according to specific gravity, are most commonly used. The degrees on Baumé'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 formulæ:—
Sp. g. = 144.3/(144.3-degrees Baumé.)
or putting the rule in words, subtract the degrees Baumé from 144.3, and divide 144.3 with the number thus obtained. For example: 32° Baumé 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 Baumé.)
or in words divide 146 by the number of degrees Baumé addedto 136. For example: ammonia of 30° Beaumé 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 Beaumé 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.