THE ALKALINE EARTHS.
LIME.
Lime is an oxide of calcium, CaO. It occurs abundantly in nature, but only in a state of combination. The carbonate (CaCO3), found as limestone, chalk, and other rocks, and as the minerals calcite and arragonite, is the most commonly occurring compound. The hydrated sulphate, gypsum (CaSO4.2H2O), is common, and is used in making "plaster of Paris." Anhydrite (CaSO4) also occurs in rock masses, and is often associated with rock salt. Phosphate of lime, in the forms of apatite, phosphorite, coprolite, &c., is largely mined. Lime is a component of most natural silicates. Calcium also occurs, combined with fluorine, in the mineral fluor (CaF2). In most of these the acid is the important part of the mineral; it is only the carbonate which is used as a source of lime.
Lime, in addition to its use in mortars and cements, is valuable as a flux in metallurgical operations, and as a base in chemical work on a large scale. A mixture of lime and magnesia is used in the manufacture of basic fire-bricks.
Carbonate of lime on ignition, especially when in contact with reducing substances, loses carbonic acid, and becomes lime. This is known as "quicklime"; on treatment with water it becomes hot, expands, and falls to a powder of "slaked lime" or calcium hydrate (CaH2O2). The hydrate is slightly soluble in water (0.1368 gram in 100 c.c.), forming an alkaline solution known as lime-water. Calcium hydrate is more generally used suspended in water as "milk of lime."
As a flux it is used either as limestone or as quicklime. Silica forms with lime a compound, calcium silicate, which is not very fusible; but when alumina and other oxides are present, as in clays and in most rocky substances, the addition of lime gives a very fusible slag.
Detection.—Calcium is detected by the reddish colour which its salts impart to the flame. It is best to moisten with hydrochloric acid (or, in the case of some silicates, to treat with ammonium fluoride) before bringing the substance into the flame. When seen through a spectroscope, it shows a large number of lines, of which a green and an orange are most intense and characteristic. Calcium is detected in solution (after removal of the metals by treatment with sulphuretted hydrogen and ammonium sulphide) by boiling with ammonium oxalate and ammonia. The lime is completely thrown down as a white precipitate. Lime is distinguished from the other alkaline earths by forming a sulphate insoluble in dilute alcohol, but completely soluble in a boiling solution of ammonium sulphate.
Lime compounds are for the most part soluble in water or in dilute hydrochloric acid. Calcium fluoride must be first converted into sulphate by evaporation in a platinum dish with sulphuric acid. Insoluble silicates are opened up by fusion with "fusion mixture," as described under Silica.
Separation.—The separation of lime is effected by evaporating with hydrochloric acid, to separate silica; and by treating with sulphuretted hydrogen, to remove the second group of metals. If the substance contains much iron, the solution is next oxidised by boiling with a little nitric acid; and the iron, alumina, &c., are removed as basic acetates. The filtrate is treated with ammonia and sulphuretted hydrogen, and allowed to settle. The filtrate from this is heated to boiling, treated with a solution of ammonium oxalate in excess, boiled for five or ten minutes, allowed to settle for half an hour, and filtered. The precipitate contains all the lime as calcium oxalate.
GRAVIMETRIC DETERMINATION.
The precipitate of calcium oxalate is washed with hot water, dried, transferred to a weighed platinum crucible, and ignited at a temperature not above incipient redness. This ignition converts the oxalate into carbonate, with evolution of carbonic oxide, which burns at the mouth of the crucible with a blue flame.[92] Generally a small quantity of the carbonate is at the same time converted into lime. To reconvert it into carbonate, moisten with a few drops of ammonic carbonate solution, and dry in a water-oven. Heat gently over a Bunsen burner, cool, and weigh. The substance is calcium carbonate (CaCO3), and contains 56 per cent. of lime (CaO). It is a white powder, and should show no alkaline reaction with moistened litmus-paper.
Where the precipitate is small, it is better to ignite strongly over the blowpipe, and weigh directly as lime. With larger quantities, and when many determinations have to be made, it is easier to make the determination volumetrically.
VOLUMETRIC METHODS.
These are carried out either by dissolving the oxalate at once in dilute sulphuric acid, and titrating with permanganate of potassium solution; or by calcining it to a mixture of lime and carbonate, and determining its neutralising power with the standard solutions of acid and alkali.
Titration with Permanganate of Potassium Solution.—This solution is made by dissolving 5.643 grams of the salt in water, and by diluting to 1 litre; 100 c.c. are equivalent to 0.5 gram of lime. The solution is standardised by titrating a quantity of oxalic acid about equivalent to the lime present in the assay; 0.5 gram of lime is equivalent to 1.125 gram of crystallised oxalic acid. The standardising may be done with iron. The standard found for iron multiplied by 0.5 gives that for lime.
The process is as follows:—The calcium oxalate (having been precipitated and washed, as in the gravimetric process) is washed through the funnel into a flask with hot dilute sulphuric acid, boiled till dissolved, diluted to 200 c.c. with water, and heated to about 80° C. The standard solution of "permanganate" is then run in, (not too quickly, and with constant shaking) until a permanent pink tinge is produced. The c.c. used multiplied by the standard, and divided by the weight of the substance taken, will give the percentage of lime.
Estimation of Lime by Alkalimetry.—The methods of determining the amount of an alkali or base by means of a standard acid solution, or, conversely, of determining an acid by means of a standard alkaline solution, are so closely related that they are best considered under one head. The same standard solution is applicable for many purposes, and, consequently, it is convenient to make it of such strength that one litre of it shall equal an equivalent in grams of any of the substances to be determined. Such solutions are termed normal. For example, a solution of hydrochloric acid (HCl = 36.5) containing 36.5 grams of real acid per litre, would be normal and of equivalent strength to a solution containing either 17 grams of ammonia (NH3 = 17) or 40 grams of sodic hydrate (NaHO = 40) per litre. It will be seen in these cases that the normal solution contains the molecular weight in grams per litre; and, if solutions of these strengths be made, it will be found that they possess equal neutralising value.
If, now, a solution containing 98 grams of sulphuric acid (H2SO4 = 98) per litre be made, it will be found to have twice the strength of the above solution, that is, 100 c.c. of the soda would only require 50 c.c. of the acid to neutralise it. The reason for this will be seen on inspecting the equations:—
NaHO + HCl = NaCl + H2O.
2NaHO + H2SO4 = Na2SO4 + 2H2O.
Acids like sulphuric acid are termed bibasic, and their equivalent is only half the molecular weight. Thus, a normal solution of sulphuric acid would contain 49 grams (98/2) of real acid per litre. Similarly, lime and most of the bases are bibasic, as may be seen from the following equations; hence their equivalent will be half the molecular weight.
2HCl + CaO = CaCl2 + H2O.
2HCl + MgO = MgCl2 + H2O.
The standard normal solution of hydrochloric acid is made by diluting 100 c.c. of the strong acid to one litre with water. This will be approximately normal. In order to determine its exact strength, weigh up 3 grams of recently ignited pure sodium carbonate or of the ignited bicarbonate. Transfer to a flask and dissolve in 200 c.c. of water; when dissolved, cool, tint faintly yellow with a few drops of a solution of methyl orange, and run in the standard "acid " from a burette till the yellow changes to a pink. Read off the number of c.c. used, and calculate to how much sodium carbonate 100 c.c. of the "acid" are equivalent. If the "acid" is strictly normal, this will be 5.3 grams. It will probably be equivalent to more than this. Now calculate how much strictly normal "acid" would be equivalent to the standard found. For example: suppose the standard found is 5.5 gram of sodium carbonate, then—
5.3 : 5.5 :: 100 : x
(where x is the quantity of normal "acid" required).
x = 103.8 c.c.
To get the "acid" of normal strength, we should then add 3.8 c.c. of water to each 100 c.c. of the standard solution remaining. Suppose there were left 930 c.c. of the approximate "acid," 35.3 c.c. of water must be added and mixed. It should then be checked by another titration with pure sodium carbonate.
The standard solution of semi-normal "alkali." The best alkali for general purposes is ammonia, but, since it is volatile (especially in strong solutions), it is best to make it of half the usual strength, or semi-normal. One litre of this will contain 8.5 grams of ammonia (NH3), and 100 c.c. of it will just neutralise 50 c.c. of the normal "acid." Take 100 c.c. of dilute ammonia and dilute with water to one litre. Run into a flask 50 c.c. of the standard "acid," tint with methyl orange, and run in from a burette the solution of ammonia till neutralised. Less than 100 c.c. will probably be used. Suppose 95 c.c. were required, there should have been 100, hence there is a deficiency of five. Then, for each 95 c.c. of standard "ammonia" left, add 5 c.c. of water, and mix well. 100 c.c. will now be equivalent to 50 c.c. of the "acid."
As an example of the application of this method, we may take the determination of lime in limestone, marble, and similar substances.
Determination of Lime in Limestone.—Weigh up 1 gram of the dried sample, and dissolve in 25 c.c. of normal acid, cool, dilute to 100 c.c., and titrate with the semi-normal solution of alkali (using methyl-orange as an indicator). Divide the c.c. of alkali used by 2, subtract from 25, and multiply by 0.028 to find the weight of lime. This method is not applicable in the presence of other carbonates or oxides, unless the weight of these substances be afterwards determined and due correction be made.
STRONTIA.
Strontia, the oxide of strontium (SrO), occurs in nature as sulphate, in the mineral celestine (SrSO4), and as carbonate in strontianite (SrCO3). It is found in small quantities in limestones, chalk, &c.
Strontia is used in sugar-refining, and for the preparation of coloured lights.
Detection.—It is detected by the crimson colour which its compounds (when moistened with hydrochloric acid) impart to the flame. The spectrum shows a large number of lines, of which a red, an orange, and a blue are most characteristic.
It resembles lime in many of its compounds, but is distinguished by the insolubility of its sulphate in a boiling solution of ammonium sulphate, and by the insolubility of its nitrate in alcohol. From baryta, which it also resembles, it is distinguished by not yielding an insoluble chromate in an acetic acid solution, by the solubility of its chloride in alcohol, and by the fact that its sulphate is converted into carbonate on boiling with a solution formed of 3 parts of potassium carbonate and 1 of potassium sulphate.
It is got into solution in the same manner as lime. The sulphate should be fused with "fusion mixture," extracted with water, and thoroughly washed. The residue will contain the strontia as carbonate, which is readily soluble in dilute hydrochloric or nitric acid.
Separation.—It is separated (after removal of the silica and metals, as described under Lime) by adding ammonia and ammonia carbonate, and allowing to stand for some hours in a warm place. In the absence of baryta or lime it is filtered off, and weighed as strontium carbonate, which contains 70.17 per cent. of strontia. It is separated from baryta by dissolving in a little hydrochloric acid, adding ammonia in excess, and then acidifying with acetic acid, and precipitating the baryta with potassium bichromate, as described under Baryta. The strontia is precipitated from the filtrate by boiling for some time with a strong solution of ammonic sulphate and a little ammonia. Fifty parts of ammonic sulphate are required for each part of strontia or lime present. The precipitate is filtered off, and washed first with a solution of ammonic sulphate, and then with alcohol. It is dried, ignited and weighed as strontium sulphate.
GRAVIMETRIC DETERMINATION.
The determination of strontia in pure solutions is best made by adding sulphuric acid in excess and alcohol in volume equal to that of the solution. Allow to stand overnight, filter, wash with dilute alcohol, dry, ignite at a red heat, and weigh as sulphate (SrSO4). This contains 56.4 per cent. of strontia (SrO); or 47.7 per cent. of strontium.
BARYTA.
Baryta, oxide of barium (BaO), commonly occurs in combination with sulphuric oxide in the mineral barytes or heavy spar (BaSO4), and in combination with carbon dioxide in witherite (BaCO3). These minerals are not unfrequently found in large quantity (associated with galena and other metallic sulphides) in lodes. Small isolated crystals of these are frequently found in mining districts. Barium is a constituent of certain mineral waters. The minerals are recognised by their high specific gravity and their crystalline form.
Compounds of barium are often used by the assayer, more especially the chloride and hydrate. The salts are, with the exception of the sulphate, generally soluble in water or hydrochloric acid. In such solutions sulphuric acid produces a white precipitate of baric sulphate, which is practically insoluble in all acids.
The dioxide (BaO2) is used for the preparation of oxygen. On strong ignition it gives up oxygen, and is converted into baryta (BaO), which, at a lower temperature, takes up oxygen from the air, re-forming the dioxide.
Detection.—Barium is detected by the green colour its salts, especially the chloride, give to the flame. This, viewed through the spectroscope, shows a complicated spectrum, of which two lines in the green are most easily recognised and characteristic. The salts of barium give no precipitate with sulphuretted hydrogen in either acid or alkaline solution, but with sulphuric acid they at once give a precipitate, which is insoluble in acetate of soda. In solutions rendered faintly acid with acetic acid, they give a yellow precipitate with bichromate of potash. These reactions are characteristic of barium.
Baryta is got into solution in the manner described under Lime; but in the case of the sulphate the substance is fused with three or four times its weight of "fusion mixture." The "melt" is extracted with water, washed, and the residue dissolved in dilute hydrochloric acid.
Separation.—The separation is thus effected:—The solution in hydrochloric acid is evaporated to dryness, re-dissolved in hot dilute hydrochloric acid, and sulphuric acid is added to the solution till no further precipitate is formed. The precipitate is filtered off, and digested with a solution of ammonium acetate or of sodium hyposulphite at 50° or 60° C. to dissolve out any lead sulphate. The residue is filtered off, washed, dried, and ignited. The ignited substance is mixed with four or five times its weight of "fusion mixture," and fused in a platinum-dish over the blowpipe for a few minutes. When cold, it is extracted with cold water, filtered, and washed. The residue is dissolved in dilute hydrochloric acid, and (if necessary) filtered. The solution contains the barium as baric chloride mixed, perhaps, with salts of strontium or lime. To separate these, ammonia is added till the solution is alkaline, and then acetic acid in slight excess. Chromate of baryta is then thrown down, by the addition of bichromate of potash, as a yellow precipitate. It is allowed to settle, filtered and washed with a solution of acetate or of nitrate of ammonia. It is dried, ignited gently, and weighed. It is BaCrO4, and contains 60.47 per cent. of baryta.
GRAVIMETRIC DETERMINATION.
The gravimetric determination of baryta, when lime and strontia are absent, is as follows:—The solution, if it contains much free acid, is nearly neutralised with ammonia, and then diluted to 100 or 200 c.c. It is heated to boiling, and dilute sulphuric acid is added till no further precipitation takes place. The precipitate is allowed to settle for a few minutes, decanted through a filter, and washed with hot water; and, afterwards, dried, transferred to a porcelain crucible, and strongly ignited in the muffle or over the blowpipe for a few minutes. It is then cooled, and weighed as sulphate of baryta (BaSO4). It contains 65.67 per cent. of baryta (BaO).
In determining the baryta in minerals which are soluble in acid, it is precipitated direct from the hydrochloric acid solution (nearly neutralised with ammonia) by means of sulphuric acid. The precipitated baric sulphate is digested with a solution of ammonic acetate; and filtered, washed, ignited, and weighed.
VOLUMETRIC DETERMINATION.
The principle and mode of working of this is the same as that given under the Sulphur Assay; but using a standard solution of sulphuric acid instead of one of barium chloride. The standard solution of sulphuric acid is made to contain 32.02 grams of sulphuric acid (H2SO4), or an equivalent of a soluble alkaline sulphate, per litre. 100 c.c. will be equal to 5 grams of baryta.
Five grams of the substance are taken, and the baryta they contain converted into carbonate (if necessary). The carbonate is dissolved in dilute hydrochloric acid. Ten grams of sodium acetate are added, and the solution, diluted to 500 c.c., is boiled, and titrated in the manner described.
Lead salts must be absent in the titration, and so must strontia and lime. Ferrous salts should be peroxidised by means of permanganate or chlorate of potash. Other salts do not interfere.
MAGNESIA.
Magnesia, the oxide of magnesium (MgO) occurs in nature in the rare mineral periclase (MgO); and hydrated, as brucite (MgH2O2). As carbonate it occurs in large quantity as magnesite (MgCO3), which is the chief source of magnesia. Mixed with carbonate of lime, it forms magnesian limestone and dolomite. It is present in larger or smaller quantity in most silicates; and the minerals, serpentine, talc, steatite and meerschaum are essentially hydrated silicates of magnesia. Soluble magnesian salts occur in many natural waters; more especially the sulphate and the chloride. Kieserite (MgSO4.H2O) occurs in quantity at Stassfurt, and is used in the manufacture of Epsom salts.
Detection.—Magnesia is best detected in the wet way. Its compounds give no colour to the flame, and the only characteristic dry reaction is its yielding a pink mass when ignited before the blowpipe (after treatment with a solution of cobalt nitrate). In solution, it is recognised by giving no precipitate with ammonia or ammonic carbonate in the presence of ammonic chloride, and by giving a white crystalline precipitate on adding sodium phosphate or arsenate to the ammoniacal solution.
Magnesia differs from the other alkaline earths by the solubility of its sulphate in water.
Magnesia is dissolved by boiling with moderately strong acids; the insoluble compounds are fused with "fusion mixture," and treated as described under Silicates.
Separation.—It is separated by evaporating the acid solution to dryness to render silica insoluble, and by taking up with dilute hydrochloric acid. The solution is freed from the second group of metals by means of sulphuretted hydrogen, and the iron, alumina, &c., are removed with ammonic chloride, ammonia, and ammonic sulphide. The somewhat diluted filtrate is treated, first, with ammonia, and then with carbonate of ammonia in slight excess. It is allowed to stand for an hour in a warm place, and then filtered. The magnesia is precipitated from the filtrate by the addition of an excess of sodium phosphate and ammonia. It is allowed to stand overnight, filtered, and washed with dilute ammonia. The precipitate contains the magnesia as ammonic-magnesic phosphate.
In cases where it is not desirable to introduce sodium salts or phosphoric acid into the assay solution, the following method is adopted. The solution (freed from the other alkaline earths by ammonium carbonate) is evaporated in a small porcelain dish with nitric acid. The residue (after removing the ammonic salts by ignition) is taken up with a little water and a few crystals of oxalic acid, transferred to a platinum dish, evaporated to dryness, and ignited. The residue is extracted with small quantities of boiling water and filtered off; while the insoluble magnesia is washed. The filtrate contains the alkalies. The residue is ignited, and weighed as magnesia. It is MgO.
GRAVIMETRIC DETERMINATION.
The solution containing the magnesia is mixed with chloride of ammonium and ammonia in excess. If a precipitate should form, more ammonic chloride is required. Add sodium phosphate solution in excess, stir and allow to stand overnight. Filter and wash the precipitate with dilute ammonia. Dry, transfer to a platinum or porcelain crucible, and ignite (finally at intense redness); cool, and weigh. The substance is magnesic pyrophosphate (Mg2P2O7), and contains 36.04 per cent. of magnesia.
VOLUMETRIC METHOD.
The magnesia having been precipitated as ammonic-magnesic phosphate, which is the usual separation, its weight can be determined volumetrically by the method of titration described under Phosphates.
The same standard solution of uranium acetate is used. Its standard for magnesia is got by multiplying the standard for phosphoric oxide by 0.5493. For example, if one hundred c.c. are equivalent to 0.5 gram of phosphoric oxide, they will be equivalent to (0.5 × .5493) 0.2746 gram of magnesia. The method of working and the conditions of the titration are the same as for the phosphate titration. The quantity of substance taken for assay must not contain more than 0.1 or 0.2 gram of magnesia. After precipitating as ammonic-magnesic phosphate with sodium phosphate, and well washing with ammonia, it is dissolved in dilute hydrochloric acid, neutralised with ammonia, and sodic acetate and acetic acid are added in the usual quantity. The solution is boiled and titrated.
EXAMINATION OF A LIMESTONE.
Silica and Insoluble Silicates.—Take one gram of the dried sample and dissolve it in 10 c.c. of dilute hydrochloric acid; filter; wash, dry, and ignite the residue.
Organic Matter.—If the residue insoluble in hydrochloric acid shows the presence of organic matter, it must be collected on a weighed filter and dried at 100°. On weighing, it gives the combined weights of organic and insoluble matter. The latter is determined by igniting and weighing again. The organic matter is calculated by difference.
Lime.—Where but little magnesia is present, this is determined by titration with standard acid. Take one gram, and dissolve it in 25 c.c. of normal hydrochloric acid. Tint with methyl-orange and titrate with semi-normal ammonia. Divide the quantity of ammonia used by 2, deduct this from 25, and multiply the remainder by 2.8. This gives the percentage of lime. Where magnesia is present, the same method is adopted, and the magnesia (which is separately determined) is afterwards deducted. The percentage of magnesia found is multiplied by 1.4, and the result is deducted from the apparent percentage of lime got by titrating.
Magnesia.—Dissolve 2 grams of the limestone in hydrochloric acid, and separate the lime with ammonia and ammonium oxalate. The filtrate is treated with sodium phosphate, and the magnesia is weighed as pyrophosphate, or titrated with uranium acetate.
Iron.—Dissolve 2 grams in hydrochloric acid, reduce, and titrate with standard permanganate of potassium solution. This gives the total iron. The ferrous iron is determined by dissolving another 2 grams in hydrochloric acid and at once titrating with the permanganate of potassium solution.
Manganese.—Dissolve 20 grams in hydrochloric acid, nearly neutralise with soda, add sodium acetate, boil, and filter. To the filtrate add bromine; boil, and determine the manganese in the precipitate. See page 300.
Phosphoric Oxide.—This is determined by dissolving the ferric acetate precipitate from the manganese separation in hydrochloric acid, adding ammonia in excess, and passing sulphuretted hydrogen. Filter and add to the filtrate "magnesia mixture." The precipitate is collected, washed with ammonia, ignited, and weighed as pyrophosphate.
THE ALKALIES.
The oxides of sodium, potassium, lithium, cæsium, and rubidium and ammonia are grouped under this head. Of these cæsia and rubidia are rare, and lithia comparatively so. They are easily distinguished by their spectra. They are characterised by the solubility of almost all their salts in water, and, consequently, are found in the solutions from which the earths and oxides of the metals have been separated by the usual group re-agents.
The solution from which the other substances have been separated is evaporated to dryness, and the product ignited to remove the ammonic salts added for the purpose of separation. The residue contains the alkali metals generally, as chlorides or sulphates. Before determining the quantities of the particular alkali metals present, it is best to convert them altogether, either into chloride or sulphate, and to take the weight of the mixed salts. It is generally more convenient to weigh them as chlorides. They are converted into this form, if none of the stronger acids are present, by simply evaporating with an excess of hydrochloric acid. Nitrates are converted into chlorides by this treatment. When sulphates or phosphates are present, the substance is dissolved in a little water, and the sulphuric or phosphoric acid precipitated with a slight excess of acetate of lead in the presence of alcohol. The solution is filtered, and the excess of lead precipitated with sulphuretted hydrogen. The filtrate from this is evaporated to dryness with an excess of hydrochloric acid, and the residue, consisting of the mixed chlorides, is gently ignited and weighed. In many cases (such as the analysis of slags and of some natural silicates where the percentage of alkalies is small) the percentage of soda and potash (which most commonly occur) need not be separately determined. It is sufficient to report the proportion of mixed alkalies; which is thus ascertained:—Dissolve the ignited and weighed chlorides in 100 c.c. of distilled water, and titrate with the standard solution of silver nitrate (using potassic chromate as indicator) in the manner described under Chlorine. The c.c. of silver nitrate used gives the weight in milligrams of the chlorine present. Multiply this by 0.775, and deduct the product from the weight of the mixed chlorides. This will give the combined weight of the alkalies (Na2O and K2O) present. For example, 0.0266 gram of mixed chlorides required on titrating 14.2 c.c. of silver nitrate, which is equivalent to 0.0142 gram of chlorine. This multiplied by 0.775 gives 0.0110 to be deducted from the weight of the mixed chlorides.
| Mixed chlorides | 0.0266 gram |
| Deduction | 0.0110 " |
| ——— | |
| Mixed alkalies | 0.0156 " |
Assuming this to have been got from 1 gram of a rock, it would amount to 1.56 per cent. of "potash and soda."
The relative proportions of the potash and soda can be ascertained from the same determination. Sodium and potassium chlorides have the following composition:—
| Sodium | 39.38 | Potassium | 52.46 |
| Chlorine | 60.62 | Chlorine | 47.54 |
| ——— | ——— | ||
| 100.00 | 100.00 |
The percentage of chlorine in the mixed chlorides is calculated. It necessarily falls somewhere between 47.5 and 60.6 per cent., and approaches the one or the other of these numbers as the proportion of the sodium or potassium preponderates. Each per cent. of chlorine in excess of 47.5 represents 7.63 per cent. of sodium chloride in the mixed chlorides. The percentage of potash and soda in the substance can be calculated in the usual way. Sodium chloride multiplied by 0.5302 gives its equivalent of soda (Na2O), and potassium chloride multiplied by 0.6317 gives its equivalent of potash (K2O).
The weight of sodium chloride in the mixed chlorides is also calculated thus:—Take the same example for illustration. Multiply the chlorine found by 2.103. This gives—
(0.0142×2.103) = 0.02987.
From the product deduct the weight of the mixed chlorides found—
| Product | 0.02987 |
| Mixed chlorides | 0.02660 |
| ——— | |
| Difference | 0.00327 |
The difference multiplied by 3.6288 gives the weight of sodium chloride in the mixture. In this case it equals 0.0118 gram. The potassium chloride is indicated by the difference between this and the weight of the mixed chlorides. It equals 0.0148 gram. We have now got—
| Sodium chloride | 0.0118 gram |
| Potassium chloride | 0.0148 " |
from 1 gram of the rock taken. Multiplying these by their factors we have (Soda = 0.0118×0.5302; Potash 0.0148×0.6317)—
| Soda | = 0.625 per cent. |
| Potash | = 0.935 " |
Concentration of the Alkalies.—With the exception of magnesia, all the other bases are separated from the alkalies in the ordinary course of work without the addition of any re-agent which cannot be removed by simple evaporation and ignition. Consequently, with substances soluble in acids, successive treatment of the solution with sulphuretted hydrogen, ammonia, ammonic sulphide, and ammonic carbonate, filtering, where necessary, will yield a filtrate containing the whole of the alkalies with ammonic salts and, perhaps, magnesia.
The filtrate is evaporated in a small porcelain dish, with the addition of nitric acid towards the finish. It is carried to dryness and ignited. The residue is taken up with a little water, treated with a few crystals of oxalic acid, and again evaporated and ignited. The alkaline salts are extracted with water, and filtered from the magnesia into a weighed platinum dish. The solution is then evaporated with an excess of hydrochloric acid, ignited at a low red heat, and weighed. The residue consists of the mixed alkaline chlorides.
For substances (such as most silicates and similar bodies) not completely decomposed by acids, Lawrence Smith's method is generally used. This is as follows:—Take from 0.5 to 1 gram of the finely powdered mineral, and mix, by rubbing in the mortar, with an equal weight of ammonium chloride. Then mix with eight times as much pure calcium carbonate, using a part of it to rinse out the mortar. Transfer to a platinum crucible, and heat gently over a Bunsen burner until the ammonic chloride is decomposed (five or ten minutes). Raise the heat to redness, and continue at this temperature for about three quarters of an hour. The crucible must be kept covered. Cool, and turn out the mass into a 4-inch evaporating dish; wash the crucible and cover with distilled water, and add the washings to the dish; dilute to 60 or 80 c.c., and heat to boiling. Filter and wash. Add to the filtrate about 1.5 gram of ammonium carbonate; evaporate to about 40 c.c., and add a little more ammonic carbonate and some ammonia. Filter into a weighed platinum dish, and evaporate to dryness. Heat gently, to drive off the ammonic chloride, and ignite to a little below redness. Cool and weigh. The residue consists of the mixed alkaline chlorides.
Separation of the Alkali-Metals from each other.—Sodium and lithium are separated from the other alkali-metals by taking advantage of the solubility of their chlorides in the presence of platinic chloride; and from one another by the formation of an almost insoluble lithic phosphate on boiling with a solution of sodium phosphate in a slightly alkaline solution. Cæsium, rubidium, and potassium yield precipitates with platinic chloride, which are somewhat soluble, and must be precipitated from concentrated solutions. Cæsium and rubidium are separated from potassium by fractional precipitation with platinum chloride. Their platino-chlorides, being less soluble than that of potassium, are precipitated first. One hundred parts of boiling water dissolve 5.18 of the potassium platino-chloride, 0.634 of the rubidium salt, and 0.377 of the corresponding cæsium compound. The separation of lithium, cæsium, and rubidium is seldom called for, owing to their rarity. The details of the separation of potassium from sodium are described under Potassium. Ammonia compounds are sharply marked off from the rest by their volatility, and it is always assumed that they have been removed by ignition; if left in the solution, they would count as potassium compounds. They will be considered under Ammonia.
SODIUM.
Sodium is the commonest of the alkali metals. It is found in nature chiefly combined with chlorine as "common salt" (NaCl). This mineral is the source from which the various compounds of sodium in use are prepared. Sodium occurs abundantly as nitrate (NaNO3) in Chili saltpetre, and as silicate in various minerals, such as albite (or soda-felspar).
It occurs as fluoride in cryolite (Na3AlF6), and as carbonate in natron, &c. Sulphates are also found. Sodium is very widely diffused, few substances being free from it.
The detection of sodium is easy and certain, owing to the strong yellow colour its salts impart to the flame; this, when viewed by the spectroscope, shows a single yellow line.[93] The extreme delicacy of this test limits its value, because of the wide diffusion of sodium salts. It is more satisfactory to separate the chloride, which may be recognised by its taste, flame coloration, fusibility, and negative action with reagents. The chloride dissolved in a few drops of water gives with potassium metantimoniate, a white precipitate of the corresponding sodium salt.
Sodium salts are dissolved out from most compounds on treatment with water or dilute acids. Insoluble silicates are decomposed and the alkali rendered soluble by Lawrence Smith's method, which has just been described. The separation of the sodium from the mixed chlorides is effected in the following way:—The chlorides are dissolved in a little water and the potassium separated as platino-chloride. The soluble sodium platino-chloride, with the excess of platinum, is boiled, mixed with sulphuric acid, evaporated to dryness, and ignited. On extracting with water, filtering, evaporating, and igniting, sodium sulphate is left, and is weighed as such.
It is more usual, and quite as satisfactory, to calculate the weight of the sodium chloride by difference from that of the mixed chlorides, by subtracting that of the potassium chloride, which is separately determined. For example, 1 gram of a rock gave—Mixed chlorides, 0.0266 gram, and 0.0486 gram of potassic platino-chloride. This last is equivalent to 0.0149 gram of potassium chloride.
| Mixed chlorides found | 0.0266 |
| Deduct potassium chloride | 0.0149 |
| ——— | |
| Leaves sodium chloride | 0.0117 |
The weight of sodium chloride found, multiplied by 0.5302, gives the weight of the soda (Na2O).
GRAVIMETRIC DETERMINATION.
The solution, which must contain no other metal than sodium, is evaporated in a weighed platinum crucible or dish. Towards the finish an excess, not too great, of sulphuric acid is added, and the evaporation is continued under a loosely fitting cover. The residue is ignited over the blowpipe, a fragment of ammonic carbonate being added towards the end, when fumes of sulphuric acid cease to be evolved. This ensures the removal of the excess of acid. The crucible is cooled in the desiccator, and weighed. The substance is sulphate of soda (Na2SO4), and contains 43.66 per cent. of soda (Na2O), or 32.38 per cent. of sodium (Na).
VOLUMETRIC METHODS.
There are various methods used for the different compounds of sodium. There is no one method of general application. Thus with "common salt" the chlorine is determined volumetrically; and the sodium, after deducting for the other impurities, is estimated by difference.
With sodic carbonate and caustic soda, a given weight of the sample is titrated with standard acid, and the equivalent of soda estimated from the alkalinity of the solution.
With sodium sulphate, a modification of the same method is used. To a solution of 3.55 grams of the salt contained in a half-litre flask, 250 c.c. of a solution of baryta water is added. The volume is made up to 500 c.c. with water. The solution is mixed and filtered. Half of the filtrate is measured off, treated with a current of carbonic acid, and then boiled. It is transferred to a half-litre flask, diluted to the mark, shaken up, and filtered. 250 c.c. of the filtrate, representing a quarter of the sample taken, is then titrated with standard acid. The standard acid is made by diluting 250 c.c. of the normal acid to 1 litre. The c.c. of acid used multiplied by 2 gives the percentage. A correction must be made to counteract the effect of impurities in the baryta as well as errors inherent in the process. This is small, and its amount is determined by an experiment with 3.55 grams of pure sodium sulphate.
EXAMINATION OF COMMON SALT.
Moisture.—Powder and weigh up 10 grams of the sample into a platinum dish. Dry in a water oven for an hour, and afterwards heat to bare redness over a Bunsen burner. Cool, and weigh. The loss gives the water.
Chlorine.—Weigh up two separate lots of 1 gram each; dissolve in 100 c.c. of water, and determine the chlorine by titrating with the standard silver nitrate solution, using chromate of potash as indicator. See Chlorine.
Insoluble Matter.—Dissolve 10 grams of the salt in water with the help of a little hydrochloric acid. Filter off the sediment, wash, ignite, and weigh. This residue is chiefly sand. Dilute the nitrate to 500 c.c.
Lime.—Take 250 c.c. of the filtrate, render ammoniacal and add ammonium oxalate; wash, dry, and ignite the precipitate. Weigh as lime (CaO).
Magnesia.—To the filtrate from the lime add phosphate of soda. Allow to stand overnight, filter, wash with dilute ammonia, dry, ignite, and weigh as pyrophosphate.
Sulphuric Oxide.—To the remaining 250 c.c. of the filtrate from the "insoluble," add an excess of barium chloride. Collect, wash, dry, ignite, and weigh the barium sulphate.
Sodium.—It is estimated by difference.
The following may be taken as an example:—
| Moisture | 0.35 |
| Insoluble matter | 0.40 |
| Lime | 0.40 |
| Magnesia | 0.05 |
| Sulphuric oxide | 0.60 |
| Chlorine | 59.60 |
| Sodium | 38.60 |
| ——— | |
| 100.00 |
POTASSIUM.
Potassium occurs in nature as chloride, in the mineral sylvine (KCl), and more abundantly combined with magnesium chloride, in earnallite (KCl.MgCl2.6H2O). It occurs as nitrate in nitre (KNO3), and as silicate in many minerals, such as orthoclase (or potash-felspar) and muscovite (or potash-mica).
Potassium compounds are detected by the characteristic violet colour they impart to the flame. The presence of sodium salts masks this tint, but the interference can be neutralised by viewing the flame through a piece of blue glass. Viewed through the spectroscope, it shows a characteristic line in the red and another in the violet. These, however, are not so easy to recognise or obtain as the sodium one. Concentrated solutions of potassium salts give a yellow crystalline precipitate with platinum chloride, and a white crystalline one with the acid tartrate of soda. For these tests the solution is best neutral. These tests are only applicable in the absence of compounds other than those of potassium and sodium.
GRAVIMETRIC DETERMINATION.
This process serves for its separation from sodium. Take 1 gram of the sample and dissolve it in an evaporating dish with 50 c.c. of water. Acidify with hydrochloric acid in quantity sufficient (if the metals are present as chlorides) to make it acid, or, if other acids are present, in at least such quantity as will provide the equivalent of chlorine. Add 3 grams of platinum, in solution as platinum chloride, and evaporate on a water-bath to a stiff paste, but not to dryness. Moisten with a few drops of platinic chloride solution without breaking up the paste by stirring. Cover with 20 c.c. of strong alcohol, and wash the crystals as much as possible by rotating the dish. Allow to settle for a few moments, and decant through a filter. Wash in the same way two or three times until the colour of the filtrate shows that the excess of the platinum chloride used is removed. Wash the precipitate on to the filter with a jet of alcohol from the wash-bottle; clean the filter-paper, using as little alcohol as possible. Dry in the water-oven for an hour. Brush the precipitate into a weighed dish, and weigh it. It is potassium platino-chloride (K2PtCl6), and contains 16.03 per cent. of potassium, or 30.56 per cent. of potassium chloride (KCl), which is equivalent to 19.3 per cent. of potash (K2O).
If the filter-paper is not free from precipitate, burn it and weigh separately. The excess of weight over that of the ash will be due to platinum and potassic chloride (Pt and 2KCl). This multiplied by 1.413 will give the weight of the potassic platino-chloride from which it was formed. It must be added to the weight of the main precipitate.
The mixed alkaline chlorides obtained in the usual course of analysis are treated in this manner; the quantity of platinum added must be about three times as much as the mixed chlorides weigh.
VOLUMETRIC METHODS.
These are the same as with soda.
Examination of Commercial Carbonate of Potash.—The impurities to be determined are moisture, silica, and insoluble matter, chlorine, sulphuric oxide, and oxide of iron. These determinations are made in the ways described under the examination of common salt.
The potassium is determined by converting it into chloride and precipitating with platinum chloride, &c., as just described.
Available Alkali.—Weigh up 23.5 grams of the sample, dissolve in water, and make up to 500 c.c. Take 50 c.c., tint with methyl orange, and titrate with the normal solution of acid. The c.c. of acid used multiplied by 2 gives the percentage of available alkali calculated as potash (K2O).
Soda.—This is calculated indirectly in the following way:—Deduct from the potassium found the quantity required for combination with the chlorine and sulphuric oxide present, and calculate the remainder to potash (K2O). The apparent surplus excess of available alkali is the measure of the soda present.
Carbon Dioxide.—The c.c. of acid used in the available alkali determination, multiplied by 2.2 and divided by 2.35, gives the percentage of carbon dioxide.
LITHIUM.
Lithia, the oxide of lithium (Li2O), occurs in quantities of 3 or 4 per cent. in various silicates, such as lepidolite (or lithia-mica), spodumene, and petalite. It also occurs as phosphate in triphyline. It is a constituent of the water of certain mineral springs. A spring at Wheal Clifford contained as much as 0.372 gram of lithium chloride per litre. In small quantities, lithia is very widely diffused.
The Detection of lithia is rendered easy by the spectroscope; its spectrum shows a red line lying about midway between the yellow sodium line and the red one of potassium. It also shows a faint yellow line. The colour of the flame (a crimson) is characteristic.
The reactions of the lithium compounds lie between those of the alkalies and of the alkaline earths. Solutions are not precipitated by tartaric acid nor by platinic chloride. The oxide is slowly soluble in water. The carbonate is not freely soluble. Lithia is completely precipitated by sodic phosphate, especially in hot alkaline solutions.
In its determination the mixed alkaline chlorides obtained in the separation of the alkalies are dissolved in water, a solution of soda is added in slight excess, and the lithia precipitated with sodic phosphate. Before filtering, it is evaporated to dryness and extracted with hot water rendered slightly ammoniacal. The residue is transferred to a filter, dried, ignited, and weighed. The precipitate is lithium phosphate (3Li2O, P2O5), and contains 38.8 per cent. of lithia. The separation of lithia from magnesia is not given by the usual authorities. Wohler recommends evaporating the solution to dryness with carbonate of soda. On extracting the residue with water, the lithia dissolves out and is determined in the filtrate. One hundred parts of water dissolve, at the ordinary temperature, 0.769 parts of lithium carbonate (Li2CO3); the basic magnesia compound is almost insoluble in the absence of carbon dioxide and ammonium salts.
CAESIUM.
The oxide of caesium, caesia (Cs2O), is found associated with lithia in lepidolite, &c., and, together with rubidium, in many mineral waters. The mineral pollux is essentially a silicate of alumina and caesia; it contains 34.0 per cent. of the latter oxide.
Caesium is best detected by the spectroscope, its spectrum being characterised by two lines in the blue and one in the red; the latter is about midway between the lithium and sodium lines.
If not detected by the spectroscope, or specially looked for, caesia would, in the ordinary course of work, be separated with the potash and weighed as potassium platino-chloride.
Caesia is separated from all the other alkalies by adding to the acid solution of the mixed chlorides a strongly acid cold solution of antimonious chloride. The acid used must be hydrochloric. The caesium is precipitated as a white crystalline precipitate (CsCl.SbCl3), which is filtered off, and washed, when cold, with strong hydrochloric acid; since it is decomposed by water or on warming. The precipitate is washed into a beaker, and treated with sulphuretted hydrogen; after filtering off the sulphide of antimony, the solution leaves, on evaporation, the caesium as chloride.
RUBIDIUM.
Rubidium occurs widely diffused in nature, but in very small quantities. It is generally associated with caesium.
It is detected by the spectroscope, which shows two violet lines and two dark red ones. Like caesium, it is precipitated with platinic chloride, and in the ordinary course of work would be weighed as potassium. It is separated from potassium by fractional precipitation with platinic chloride. Rubidium platino-chloride is much less soluble than the potassium salt.
AMMONIUM.
It is usual to look upon the salts of ammonia as containing a compound radical (NH4 = Am), which resembles in many respects the metals of the alkalies. Ammonium occurs in nature as chloride in sal ammoniac (AmCl), as sulphate in mascagnine (Am2SO4), as phosphate in struvite (AmMgPO4.12H2O). Minerals containing ammonium are rare, and are chiefly found either in volcanic districts or associated with guano. Ammonia and ammonium sulphide occur in the waters of certain Tuscan lagoons, which are largely worked for the boracic acid they contain. The crude boracic acid from this source contains from 5 to 10 per cent. of ammonium salts. It is from these that the purer forms of ammonium compounds of commerce known as "from volcanic ammonia" are derived. But the bulk of the ammonia of commerce is prepared from the ammoniacal liquors obtained as bye-products in the working of certain forms of blast furnaces and coke ovens, and more especially in gas-making.
Ammonia hardly comes within the objects of assaying; but it is largely used in the laboratory, and the assayer is not unfrequently called on to determine it. Ammonium salts are mostly soluble in water. In strong solutions they give a yellow precipitate of ammonium platino-chloride on the addition of chloride of platinum; and with the acid tartrate of soda yield a white precipitate of hydric ammonic tartrate. These reactions are similar to those produced with potassium compounds.
Heated with a base, such as lime or sodic hydrate, ammonium salts are decomposed, yielding ammonia gas (NH3), which is readily soluble in water. The solution of this substance is known as ammonic hydrate or "ammonia."
They are volatilised on ignition; either with, or without, decomposition according to the acid present. This fact is of importance in analytical work; since it allows of the use of alkaline solutions and reagents which leave nothing behind on heating. It must be remembered, however, that, although ammonic chloride is volatile, it cannot be volatilised in the presence of substances which form volatile chlorides without loss of the latter. For example: ferric oxide and alumina are thus lost, volatilising as chlorides; and there are some other compounds (notably ammonic magnesic arsenate) which on heating to redness suffer reduction. The presence of ammonic chloride in such cases must be avoided.
Detection.—Compounds of ammonium are detected by their evolving ammonia when mixed or heated with any of the stronger bases. The ammonia is recognised by its odour, by its alkaline reaction with litmus paper, and by yielding white fumes, when brought in contact with fuming acid. In consequence of the use of ammonium salts and ammonia as reagents, it is necessary to make a special test for and determination of ammonium.[94] In the ordinary course of work it will be "lost on ignition." The determination presents little difficulty, and is based on the method used for its detection.
Solution and Separation.—Although ammonium salts are soluble in water, there is no necessity for dissolving them. The compound containing the ammonia is boiled with an alkaline solution; and the liberated ammonia condensed and collected. The substance is weighed out into a flask of about 200 c.c. capacity. The flask is closed with a rubber cork perforated to carry a 20 c.c. pipette and a bulb exit tube. The latter is connected with a receiver, which is a small flask containing dilute hydrochloric acid (fig. 61). The flask containing the substance is corked, and the greater part of the soda solution is run in from the pipette. The solution is then boiled. The ammonia volatilises, and is carried over into the hydrochloric acid, with which it combines to form ammonic chloride. The distillation is carried on gently until the bulk of the liquid is driven over. The ammonia in the receiver will be mixed only with the excess of hydrochloric acid. This separation is used in all determinations.
GRAVIMETRIC DETERMINATION.
The contents of the flask are transferred to a weighed platinum dish, and evaporated on the water-bath. It is dried until the weight is constant. The chloride of ammonium remains as a white mass which, after cooling in a desiccator, is weighed. It contains 33.72 per cent. of ammonium (NH4), or 31.85 per cent. of ammonia (NH3). On heating over the Bunsen burner it is completely volatilised, leaving no residue.
VOLUMETRIC DETERMINATION.
Weigh up 1.7 gram of the substance and place it in the flask. Measure off 50 c.c. of the normal solution of acid, place them in the receiver, and dilute with an equal volume of water. Run in through the pipette (by opening the clip) 20 c.c. of a strong solution of soda, boil until the ammonia has passed over, and then aspirate a current of air through the apparatus. Disconnect the receiver, and tint its contents with methyl orange. Titrate the residual acid with a semi-normal solution of alkali. Divide the c.c. of the "alkali" solution used by 2, and deduct from the 50 c.c. The difference will give the number of c.c. of the normal acid solution neutralised by the ammonia distilled over. Each c.c. of "acid" so neutralised, represents 1 per cent. of ammonia in the sample. If the results are to be reported as ammonium, 1.8 gram of the sample is taken instead of 1.7 gram.
COLORIMETRIC DETERMINATION.
This is effected by means of "Nessler's" reagent, which strikes a brown colour with traces of ammonia, even with a few hundredths of a milligram in 100 c.c. of liquid. With larger quantities of ammonia the reagent gives a precipitate. This reagent is a strongly alkaline solution of potassic mercuric iodide; and is thus made:—
Nessler's solution: Dissolve 17 grams of mercuric chloride in 300 c.c. of water; and add the solution to one of 35 grams of potassium iodide in 100 c.c. of water until a permanent precipitate is produced. Both solutions must be cold. Then make up to a litre by adding a 20 per cent. solution of potash. Add more of the mercuric chloride (a little at a time) until a permanent precipitate is again formed. Allow to settle, decant, and use the clear liquor. Four or five c.c. are used for each 100 c.c. of liquid to be tested.
A Standard Solution of Ammonia is made by dissolving 0.315 gram of ammonic chloride in water, and diluting to 100 c.c. Ten c.c. of this are taken and diluted to 1 litre. One c.c. contains 0.01 milligram of ammonia (NH3).
In working, the solution containing the ammonia is diluted to a definite volume, and to such an extent that 50 c.c. of it shall not contain more than 0.02 or 0.03 milligram of ammonia. Fifty c.c. of it are transferred to a Nessler glass and mixed with 2 c.c. of Nessler's reagent. The colour is noted, and an estimate made as to the amount of ammonia it indicates. A measured quantity of the standard ammonia, judged to contain about as much ammonia as that in the assay, is then put into another Nessler glass. It is diluted to 50 c.c. with water, and mixed with 2 c.c. of "Nessler." After standing a minute or two, the colours in the two glasses are compared. If the tints are equal, the assay is finished; but if the standard is weaker or stronger than the assay, another standard, containing more or less ammonia, as the case may be, must be prepared and compared with the assay. Two such experiments will generally be sufficient; but, if not, a third must be made. The addition of more standard ammonia to the solution to which the "Nessler" has already been added does not give a satisfactory result.
When the ammonia in 50 c.c. has been determined, that in the whole solution is ascertained by a suitable multiplication. By 10, for example, if the bulk was 500 c.c., or by 20 if it was a litre.
Distilled water is used throughout. It must be free from ammonia; and is best prepared by distilling an ammonia-free spring water.