FOOTNOTES:

FOOTNOTES:[78]MnO2+ 4HCl = MnCl2+ Cl2+ 2H2O.[79]Provided a sufficiency of ammonic chloride is present.[80]With some silicates, &c., a preliminary fusion with sodium carbonate will be necessary.[81]Instead of sodium acetate, ammonium succinate can be used.[82]Journ. Soc. Chem. Industry, vol. x. p. 333.[83]MnO2+ 2FeSO4+ 2H2SO4= Fe2(SO4)3+ MnSO4+ 2H2O.[84]If the ore is very rich, a smaller quantity (0.75 or 1.5 gram) must be taken; otherwise the iron will be insufficient.[85]MnO2+ 4HCl = MnCl2+ 2H2O + Cl2.Cl2+ 2KI = 2KCl + I2.[86]Iodine probably lost by volatilisation.[87]Obtained as a brown powder by digesting red lead with nitric acid and filtering.[88]The water for dilution and the dilute sulphuric acid used for washing should be previously tested, to see they have no reducing action, with dilute permanganate of potassium solution.[89]Arnold and Hardy,Chemical News, vol. lvii. p. 153.

[78]MnO2+ 4HCl = MnCl2+ Cl2+ 2H2O.

[79]Provided a sufficiency of ammonic chloride is present.

[80]With some silicates, &c., a preliminary fusion with sodium carbonate will be necessary.

[81]Instead of sodium acetate, ammonium succinate can be used.

[82]Journ. Soc. Chem. Industry, vol. x. p. 333.

[83]MnO2+ 2FeSO4+ 2H2SO4= Fe2(SO4)3+ MnSO4+ 2H2O.

[84]If the ore is very rich, a smaller quantity (0.75 or 1.5 gram) must be taken; otherwise the iron will be insufficient.

[85]MnO2+ 4HCl = MnCl2+ 2H2O + Cl2.Cl2+ 2KI = 2KCl + I2.

[86]Iodine probably lost by volatilisation.

[87]Obtained as a brown powder by digesting red lead with nitric acid and filtering.

[88]The water for dilution and the dilute sulphuric acid used for washing should be previously tested, to see they have no reducing action, with dilute permanganate of potassium solution.

[89]Arnold and Hardy,Chemical News, vol. lvii. p. 153.

Alumina, the oxide of aluminium (Al2O3), is found in nature fairly pure in the mineral corundum; transparent and coloured varieties of which form the gems sapphire and ruby. A coarser compact variety contaminated with oxide of iron constitutes emery. Compounded with silica, alumina forms the base of clays and many rock-forming minerals. China clay (or kaolin) is used as a source of alumina. Bauxite, hydrated alumina, is also used for the same purpose—that is, for the preparation of sulphate of alumina. The mineral cryolite is a fluoride of aluminium and sodium.

Corundum is characterised by a high specific gravity (4.0) and extreme hardness. By these it is distinguished from felspar and similar minerals, which it somewhat resembles in general appearance.

Aluminium is used for a variety of small purposes: it is white, light, and very tenacious; but owing to the difficulty of its reduction it is expensive.

Aluminium forms one series of salts which closely resemble those of ferric iron. It forms an interesting series of double sulphates, known as the alums. Common potash alum is Al2(SO4)3,K2SO4,24H2O.

Detection.—Alumina is not precipitated from its acid solution by sulphuretted hydrogen, but it is thrown down by ammonia (with the other earths) as a white hydrate, soluble in soda and insoluble in ammonic carbonate. Filtered off and ignited, it assumes, after treatment with nitrate of cobalt before the blowpipe, a blue colour which is characteristic. With natural compounds containing metallic oxides this colour is masked. It is more satisfactory to make a separation in the wet way and to test the ignited oxide.

Separation and Solution.—If the substance is insoluble in hydrochloric acid it is finely powdered and fused with "fusion mixture" with the help, in the case of corundum (which is very refractory) of a little caustic soda or potash. The method ofworking is the same as that described for opening up silicates. See underSilica. Corundum cannot be powdered in Wedgwood, or even agate, mortars; since it rapidly wears these away and becomes contaminated with their powder. It is best to use a hard steel mortar and to extract the metallic particles from the bruised sample with a magnet or dilute acid.

When the substance has been completely attacked and dissolved, it is evaporated to dryness with an excess of hydrochloric acid on the water-bath to render any silica present insoluble. The residue is extracted with hydrochloric acid and freed from the second group of metals by means of sulphuretted hydrogen. The filtrate from this (after removing the sulphuretted hydrogen by boiling) is nearly neutralised, and treated with 8 or 10 grams of hyposulphite of soda[90]in solution. It is then boiled till the sulphurous oxide is driven off. The precipitate is filtered off, ignited, and weighed as alumina.

It is sometimes more convenient to proceed as follows:—After boiling off the sulphuretted hydrogen peroxidise the iron with a little nitric acid, add a solution of ammonic chloride, and then ammonia in very slight excess; boil, filter, wash, ignite, and weigh the oxides. These generally consist of ferric oxide and alumina. It is a common practice to determine the iron, calculate it to ferric oxide, and so to estimate the alumina indirectly. This may be done either by igniting in a current of hydrogen and estimating the iron by the weight of oxygen lost; or, by dissolving with sulphuric and hydrochloric acids, and determining the iron volumetrically. It should be borne in mind that these oxides will also contain any phosphoric oxide that happened to be in the mineral.

In general analyses of samples containing alumina, it may be contained in both the soluble and insoluble portions. In these cases it is better to fuse the sample with "fusion mixture" before treatment with acids. The alumina in the fused mass will exist in a state soluble in acids.

Solutions containing alumina free from the other metals are diluted to a convenient bulk and heated nearly to boiling. Add chloride of ammonium, and then ammonia in slight excess; boil, allow to settle, filter, and wash with hot water. Dry the precipitate, and ignite in a platinum or porcelain crucible at the strongest heat. Cool, and weigh. The substanceis alumina, Al2O3, which contains 52.94 per cent. of aluminium. It is only in special cases, such as the analysis of metals and alloys, that it is reported as aluminium. The percentage of alumina is generally given.

Ignited alumina is difficultly soluble in acids; it is not reduced by hydrogen at a red heat. Ignited with ammonium chloride portions are volatilised.

Direct Determination of Alumina in the Presence of Iron.—The iron and alumina are precipitated as hydrates by ammonia. The precipitate is dissolved in hydrochloric acid and the iron reduced to the ferrous state. It is then added to a hot solution of potash or soda. The solution is boiled till the precipitate settles readily, filtered, and washed with hot water. The alumina is contained in the filtrate, which is acidified with hydrochloric acid and the alumina precipitated therefrom as hydrate with ammonia, as just described.

Determination of Alumina in the Presence of Phosphates and Iron.—For details, see a paper by R.T. Thomson in the "Journal of the Society of Chemical Industry," v. p. 152. The principles of the method are as follows:—If the substance does not already contain sufficient phosphoric oxide to saturate the alumina, some phosphate is added. The iron is reduced to the ferrous state and phosphate of alumina precipitated in an acetic acid solution. It is purified by reprecipitation, ignited, and weighed as phosphate (Al2O3,P2O5), which contains 41.8 per cent. of alumina, Al2O3.

Moisture.—Take 5 grams of the carefully-prepared sample and dry in the water-oven till the weight is constant.

Loss on Ignition.—Weigh up 2 grams of the sample used for the moisture determination, and ignite in a platinum-crucible to redness, cool, and weigh.

Silica and Insoluble Silicates.—Weigh up another 2 grams of the dried sample, and place them in a platinum dish; moisten with water, and cover with 20 c.c. of sulphuric acid. Evaporate and heat gently to drive off the greater portion of the free acid. Allow to cool; and repeat the operation. Extract by boiling with dilute hydrochloric acid, filter, wash, dry, ignite, and weigh. The quantity of insoluble silicates is determined by dissolving out the separated silica with a strong boiling solution of sodium carbonate. The residue (washed, dried, and ignited) is weighed, and reported as "sand."

Alumina and Ferrous Oxide.—To the filtrate from the silica add "soda" solution till nearly neutral, and then sodiumacetate. Boil and filter off the precipitate. Reserve the filtrate. Dissolve the precipitate in hydrochloric acid, and dilute to exactly 200 c.c. Divide into two parts of 100 c.c. each. In one determine the iron by reducing and titrating in the way described under volumetric iron. Calculate the percentage as ferrous oxide, unless there are reasons to the contrary, also calculate its weight as ferric oxide. To the other portion add ammonia in slight excess, and boil. Filter, wash with hot water, dry, ignite, and weigh as mixed alumina and ferric oxide. The weight of the ferric oxide has already been determined in the first portion: deduct it, and the difference is the weight of alumina.

Lime.—To the reserved filtrate, concentrated by evaporation, add ammonium oxalate and ammonia; boil, filter, ignite strongly, and weigh as lime.

Magnesiais separated from the filtrate by adding sodium phosphate. It is weighed as magnesium pyrophosphate.

Potash and Soda.—These are determined in a fresh portion of the sample by Lawrence Smith's method, as described on page 333.

This is an oxide of thorium, ThO2. It is only found in a few rare minerals. It is a heavy oxide, having, when strongly ignited, a specific gravity of 9.2. In the ordinary course of analysis it will be separated and weighed as alumina. It is separated from this and other earths by the following method. The solution in hydrochloric acid is nearly neutralised and then boiled with sodium hyposulphite. The thoria will be in the precipitate. It is dissolved, and the solution heated with ammonium oxalate in excess. The precipitate is thorium oxalate, which is washed with hot water, dried, and ignited. It is then weighed as thoria, ThO2. Thoria which has been ignited is not readily soluble in acids.

The oxide of zirconium, ZrO2, is found in the mineral zircon, a silicate of zirconia, ZrSiO4. When heated intensely it becomes very luminous, and is used on this account for incandescent lights.

In the ordinary course it is thrown down by ammonia with the other earths, from which it is thus separated:—The hydrates precipitated in the cold, and washed with cold water, are dissolvedin hydrochloric acid, nearly neutralised with soda, and precipitated by boiling with hyposulphite of soda. Dissolve; and from the hydrochloric acid solution precipitate the thoria (if any) with ammonium oxalate. To the filtrate add carbonate of ammonia, which will precipitate any titanium present. The zirconia will be in solution, and is recovered by precipitating with potassium sulphate, or by evaporating the solution and igniting. It is separated from alumina by taking advantage of its insolubility in potassic hydrate.

It is estimated in zircons in the following way:—The powdered substance is fused with bisulphate of potash, and extracted with dilute sulphuric acid. The residue is fused with caustic soda and extracted with water. The portion not dissolved, consisting of zirconate of soda, is dissolved in hydrochloric acid. The solution is diluted, filtered if necessary, and treated with ammonia in excess. The precipitate is filtered off, washed with hot water, dried, ignited, and weighed as zirconia, ZrO2. This is a white powder, which is insoluble in acids; even in hydrofluoric acid it is only slightly attacked.

Cerium occurs as silicate (together with the oxides of lanthanum, didymium, iron and calcium) in the mineral cerite, which is its chief source. It also occurs as phosphate in monazite, and as fluoride in fluocerite. The oxalate is used in medicine. Cerium forms two classes of salts corresponding to the oxides, cerous oxide (Ce2O3) and ceric oxide (CeO2). Compounds of cerium with volatile acids yield dioxide on ignition; and this, on solution in hydrochloric acid, yields cerous chloride and chlorine.

In the ordinary course cerium is thrown down along with alumina and the other earths by ammonia. It is separated by dissolving the hydrates in hydrochloric acid, and oxidizing with chlorine water. On treating with oxalic acid, cerium, lanthanum, and didymium are precipitated as oxalates, which on ignition are converted into oxides. These are soluble in acids. Their solution in hydrochloric acid is nearly neutralised; acetate of soda is then added, and an excess of sodium hypochlorite. On boiling, the cerium is precipitated as dioxide, which is filtered off, ignited, and weighed.

Cerium is detected by giving with borax a bead which is yellow in the oxidising, and colourless in the reducing flame. Traces of cerium compounds boiled with dioxide of lead and nitric acid will give a yellow solution.

occur together with cerium in cerite, and are separated with that metal as oxalates, as described underCerium.

Didymium salts have a rose or violet colour, and impart (when in sufficient quantity) the same colour to the borax bead. Solutions have a characteristic absorption-spectrum.

The separation of lanthanum and didymium in the solution from which the cerium has been precipitated is effected by precipitating them together as oxalates, igniting, and dissolving in dilute nitric acid. This solution is then evaporated to dryness and ignited, for a few minutes, just below redness. A subnitrate of didymium is formed, and remains as an insoluble residue on extracting with hot water. The separated salts are treated with ammonia and ignited, and weighed as oxides (La2O3and Di2O3).

Yttria is found in gadolinite and some other rare minerals. It is precipitated along with the other earths by ammonia. It is distinguished by the insolubility of its hydrate in potash, by the insolubility of its oxalate in oxalic acid, and by not being precipitated by hyposulphite of soda or potassium sulphate. Further, it is precipitated by potash in the presence of tartaric acid as an insoluble tartrate. This reaction distinguishes the members of the yttria group from most of the other earths. The other members of the group closely resemble it, and amongst them are erbia, terbia, ytterbia, scandia, &c.

The oxide of beryllium, BeO (also known as glucina), occurs in nature mainly as silicate. Beryl, the green transparent variety of which is the emerald, is the best known of these. It is a silicate of alumina and beryllia.[91]Some other minerals in which it occurs are phenakite, euclase, and chrysoberyl.

In the ordinary course of analysis, beryllia will be precipitated with alumina, &c., by ammonic hydrate. It is distinguished by the solubility of its hydrate in ammonic carbonate, by not being precipitated by boiling with sodium hyposulphite, and by not being precipitated by ammonic sulphide from an ammonic carbonate solution.

The analysis of silicates containing beryllia is thus effected.The finely powdered substance is fused with twice its weight of potassium carbonate; and the "melt" is extracted with water, and evaporated with a slight excess of sulphuric acid to render the silica insoluble. Treat with water, filter, and evaporate the filtrate until a crust is formed. Potash alum crystallises out. The liquor is poured off into a warm strong solution of ammonium carbonate. Ferric hydrate and alumina will be precipitated. They are filtered off, re-dissolved, and again precipitated in ammonic carbonate solution; the combined filtrates are boiled for some time, and acidified slightly with hydrochloric acid. The carbon dioxide is boiled off, and the beryllia is then precipitated as hydrate with ammonia. The hydrate is washed with hot water, dried, ignited, and weighed as beryllia, BeO.

Beryllia has a specific gravity of 3.08. It is white, infusible, and insoluble in water. After ignition, it is insoluble in acids, except sulphuric, but is rendered soluble by fusion with alkalies.

Beryllia, in a solution of carbonate of ammonia, is precipitated as carbonate on boiling in proportion as the carbonate of ammonia is volatilised. The hydrate is dissolved by a boiling solution of ammonic chloride, ammonia being evolved.

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. Thisis 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 underSilica.

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.

The precipitate of calcium oxalate is washed with hot water, dried, transferred to a weighed platinum crucible, and ignited ata 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.

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 thatthey 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 termednormal. 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 acidis 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 standardfound. For example: suppose the standard found is 5.5 gram of sodium carbonate, then—

5.3 : 5.5 :: 100 :x(wherexis 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, orsemi-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, 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 itscompounds (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 underLime) 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 underBaryta. 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.

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, 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 underLime; 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 blowpipefor 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.

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.

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, 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 underSilicates.

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 isadopted. 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.

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.

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 underPhosphates.

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.

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 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 underChlorine. 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.

Back to Index Next