XIV. EARTHS, ALKALINE EARTHS, ALKALIES
ALUMINA.
Alumina, the oxide of aluminium (Al_{2}O_{3}), 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 Al_{2}(SO_{4}){3},K{2}SO_{4},24H_{2}O.
~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 of working is the same as that described for opening up silicates. See under Silica. 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.
GRAVIMETRIC DETERMINATION.
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 substance is alumina, Al_{2}O_{3}, 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 (Al_{2}O_{3},P_{2}O_{5}), which contains 41.8 per cent. of alumina, Al_{2}O_{3}.
EXAMINATION OF CLAYS.
~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 sodium acetate. 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.
~Magnesia~ is 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.
THORIA.
This is an oxide of thorium, ThO_{2}. 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, ThO_{2}. Thoria which has been ignited is not readily soluble in acids.
ZIRCONIA.
The oxide of zirconium, ZrO_{2}, is found in the mineral zircon, a silicate of zirconia, ZrSiO_{4}. 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 dissolved in 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, ZrO_{2}. This is a white powder, which is insoluble in acids; even in hydrofluoric acid it is only slightly attacked.
CERIUM.
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 (Ce_{2}O_{3}) and ceric oxide (CeO_{2}). 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.
LANTHANUM AND DIDYMIUM
occur together with cerium in cerite, and are separated with that metal as oxalates, as described under Cerium.
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 (La_{2}O_{3} and Di_{2}O_{3}).
YTTRIA.
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.
BERYLLIA.
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.
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 (CaCO_{3}), found as limestone, chalk, and other rocks, and as the minerals calcite and arragonite, is the most commonly occurring compound. The hydrated sulphate, gypsum (CaSO_{4}.2H_{2}O), is common, and is used in making "plaster of Paris." Anhydrite (CaSO_{4}) 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 (CaF_{2}). 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 (CaH_{2}O_{2}). 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 (CaCO_{3}), 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 (NH_{3} = 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 (H_{2}SO_{4} = 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 + H_{2}O. 2NaHO + H_{2}SO_{4} = Na_{2}SO_{4} + 2H_{2}O.
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 = CaCl_{2} + H_{2}O. 2HCl + MgO = MgCl_{2} + H_{2}O.
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 (NH_{3}), 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 (SrSO_{4}), and as carbonate in strontianite (SrCO_{3}). 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 (SrSO_{4}). 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 (BaSO_{4}), and in combination with carbon dioxide in witherite (BaCO_{3}). 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 (BaO_{2}) 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 BaCrO_{4}, 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 (BaSO_{4}). 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 (H_{2}SO_{4}), 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 (MgH_{2}O_{2}). As carbonate it occurs in large quantity as magnesite (MgCO_{3}), 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 (MgSO_{4}.H_{2}O) 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 (Mg_{2}P_{2}O_{7}), 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 (Na_{2}O and K_{2}O) 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 (Na_{2}O), and potassium chloride multiplied by 0.6317 gives its equivalent of potash (K_{2}O).
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 (NaNO_{3}) in Chili saltpetre, and as silicate in various minerals, such as albite (or soda-felspar).
It occurs as fluoride in cryolite (Na_{3}AlF_{6}), 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 (Na_{2}O).
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 (Na_{2}SO_{4}), and contains 43.66 per cent. of soda (Na_{2}O), 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.MgCl_{2}.6H_{2}O). It occurs as nitrate in nitre (KNO_{3}), 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 (K_{2}PtCl_{6}), 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 (K_{2}O).
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 (K_{2}O).
~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 (K_{2}O). 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 (Li_{2}O), 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 (3Li_{2}O, P_{2}O_{5}), 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 (Li_{2}CO_{3}); the basic magnesia compound is almost insoluble in the absence of carbon dioxide and ammonium salts.
CAESIUM.
The oxide of caesium, caesia (Cs_{2}O), 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.SbCl_{3}), 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 (NH_{4} = 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 (Am_{2}SO_{4}), as phosphate in struvite (AmMgPO_{4}.12H_{2}O). 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 (NH_{3}), 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 (NH_{4}), or 31.85 per cent. of ammonia (NH_{3}). 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 (NH_{3}).
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.
FOOTNOTES:
[90] Al_{2}Cl_{6} + 3Na_{2}S_{2}O_{3} + 3H_{2}O = Al_{2}(HO){6} + 6NaCl + 3S + 3SO{2}
[91] 3BeO,Al_{2}O_{3},6SiO_{2}
[92] CaC_{2}O_{4} = CaCO_{3}+CO.
[93] Resolved into two with a powerful spectroscope.
[94] Ammonium compounds are frequently produced when dissolving metals in nitric acid; or when nitrates are heated in the presence of the metals.
XV. OXYGEN AND OXIDES.--THE HALOGENS
OXYGEN.
Oxygen occurs in nature in the free state, forming 23 per cent. by weight, or 21 per cent. by volume of the atmosphere; but, since it is a gas, its presence is easily overlooked and its importance underestimated. Except in the examination of furnace-gases, &c., the assayer is not often called upon to determine its quantity, but it forms one of his most useful reagents, and there are many cases where he cannot afford to disregard its presence. It occurs not only in the air, but also dissolved in water; ordinary waters containing on an average 0.00085 per cent. by weight, or 0.85 parts per 100,000.
Chemically, it is characterised by its power of combining, especially at high temperatures, with the other elements, forming an important class of compounds called oxides. This combination, when rapid, is accompanied by the evolution of light and heat; hence oxygen is generally called the supporter of combustion. This property is taken advantage of in the operation of calcining, scorifying, cupelling, &c. The importance of a free access of air in all such work is seen when it is remembered that 1 litre of air contains 0.2975 gram of oxygen, and this quantity will only oxidise 0.1115 gram of carbon, 0.2975 gram of sulphur, or 3.849 grams of lead.
Oxidation takes place at the ordinary temperature with many substances. Examples of such action are seen in the weathering of pyrites, rusting of iron, and (in the assay office) the weakening of solutions of many reducing agents.
For methods of determining the percentage of oxygen in gases, for technical purposes, the student is referred to Winkler & Lunge's "Technical Gas Analysis."
OXIDES.
Oxides are abundant in nature, almost all the commonly occurring bodies being oxidised. Water (H_{2}O) contains 88.8 per cent. of oxygen; silica, lime, alumina, magnesia, and the other earths are oxides, and the oxides of the heavier metals are in many cases important ores; as, for example, cassiterite (SnO_{2}), hæmatite (Fe_{2}O_{3}), magnetite (Fe_{3}O_{4}), and pyrolusite (MnO_{2}). In fact, the last-named mineral owes its value to the excess of oxygen it contains, and may be regarded as an ore of oxygen rather than of manganese.
Most of the metals, when heated to redness in contact with air, lose their metallic lustre and become coated with, or (if the heating be prolonged) altogether converted into, oxide. This oxide was formerly termed a "calx," and has long been known to weigh more than the metal from which it was obtained. For example, one part by weight of tin becomes, on calcining, 1.271 parts of oxide (putty powder). The student will do well to try the following experiments:--Take 20 grams of tin and heat them in a muffle on a scorifier, scraping back the dross as it forms, and continuing the operation until the whole of the metal is burnt to a white powder and ceases to increase in weight.[95] Take care to avoid loss, and, when cold, weigh the oxide formed. The oxide should weigh 25.42 grams, which increase in weight is due to the oxygen absorbed from the air and combined with the metal. It can be calculated from this experiment (if there has been no loss) that oxide of tin contains 21.33 per cent. of oxygen and 78.67 per cent. of tin. Oxidation is performed with greater convenience by wet methods, using reagents, such as nitric acid, which contain a large proportion of oxygen loosely held. Such reagents are termed oxidising agents. Besides nitric acid, permanganate of potash, bichromate of potash, and peroxide of hydrogen are largely used for this purpose. One c.c. of nitric acid contains as much oxygen as 2.56 litres of air, and the greater part of this is available for oxidising purposes. Try the following experiment:--Take 2 grams of tin and cover in a weighed Berlin dish with 20 c.c. of dilute nitric acid, heat till decomposed, evaporate to dryness, ignite, and weigh. The 2 grams of tin should yield 2.542 grams of oxide. The increase in weight will be proportionally the same as in the previous experiment by calcination, and is due to oxygen, which in this case has been derived from the nitric acid.
The percentage of oxygen in this oxide of tin (or in any of the oxides of the heavier metals) may be directly determined by heating such oxides in a current of hydrogen, and collecting and weighing the water formed.
It is found by experiment that 88.86 parts by weight of oxygen, combining with 11.14 parts of hydrogen, form 100 parts of water; so that from the weight of water formed it is easy to calculate the amount of oxygen the oxide contained.
Take 1 gram of the dried and powdered oxide and place it in a warm dry combustion tube. Place the tube in a furnace, and connect at one end with a hydrogen apparatus provided with a sulphuric acid bulb for drying the gas, and at the other with a weighed sulphuric acid tube for collecting the water formed. The apparatus required is shown in fig. 62. Pass hydrogen through the apparatus, and, when the air has been cleared out, light the furnace. Continue the heat and current of hydrogen for half an hour (or longer, if necessary). Allow to cool. Draw a current of dry air through the weighed tube. Weigh. The increase in weight gives the amount of water formed, and this, multiplied by 0.8886, gives the weight of the oxygen. The percentage of oxygen thus determined should be compared with that got by the oxidation of the metal. It will be practically the same. The following results can be taken as examples:--
Twenty grams of tin, calcined as described, gave 25.37 grams of oxide.
Two grams of tin, oxidised with nitric acid and ignited, gave 2.551 grams of oxide.
One gram of the oxide of tin, on reduction in a current of hydrogen, gave 0.2360 gram of water (equivalent to 0.2098 gram of oxygen), and left 0.7900 gram of metal.
Ten grams of ferrous sulphate gave, on strong ignition, 2.898 grams of ferric oxide (Fe_{2}O_{3})[96] instead of 2.877.
The student should similarly determine the percentage of oxygen in oxides of copper and iron. The former oxide may be prepared by dissolving 5 grams of copper in 50 c.c. of dilute nitric acid, evaporating to dryness, and strongly igniting the residue. The oxide of iron may be made by weighing up 10 grams of powdered ferrous sulphate (= to 2.014 grams of iron) and heating, at first gently, to drive off the water, and then at a red heat, until completely decomposed. The weight of oxide, in each case, should be determined; and the percentage of oxygen calculated. Compare the figures arrived at with those calculated from the formula of the oxides, CuO and Fe_{2}O_{3}.
It would be found in a more extended series of experiments that the same metal will, under certain conditions, form two or more oxides differing among themselves in the amount of oxygen they contain. These oxides are distinguished from one another by such names as "higher" and "lower oxides," "peroxides," "protoxides," "dioxides," &c.
The oxides may be conveniently classified under three heads:--
(1) Those that are reduced to metal by heat alone, such as the oxides of mercury, silver, platinum, gold, &c.;
(2) Those which are reduced by hydrogen at a red heat, which includes the oxides of the heavy metals;
(3) Those which are not reduced by these means, good examples of which are silica, alumina, the alkalies, and the alkaline earths.
Another important classification is into acid, basic and neutral oxides. The oxides of the non-metallic elements, such as sulphur, carbon, phosphorus, &c., are, as a rule, acid; and the more oxygen they contain, the more distinctly acid they are. The oxides of the metals are nearly all basic; and, as a rule, the less oxygen they contain, the more distinctly basic they are.
The basic oxides, which are soluble in acids, give rise to the formation of salts when dissolved therein. During the solution, water is formed, but no gas is evolved. The oxide dissolved in each case neutralizes an equivalent of the acid used for solution.[97] The basic properties of many of these can be taken advantage of for their determination. This is done in the case of soda, potash, lime, &c., by finding the quantity of acid required to neutralize a given weight of the substance.
There are some oxides which, under certain conditions, are acid to one substance (a stronger base) and basic to another (a stronger acid). For example, the oxides of lead and of tin, as also alumina, dissolve in caustic soda, acting as acids; whilst, on the other hand, they combine with sulphuric or hydrochloric acid, playing the part of bases.
The oxides known as "earths," when ignited, are many of them insoluble in acids, although easily dissolved before ignition.
It is common in complete analyses of minerals to meet with cases in which the sum total of the elements found falls short of the amount of ore taken; and here oxygen must be looked for. For example, this occurs in the case of a mixture of pyrites with oxide of iron, or in a mixture of sulphides and sulphates. The state in which the elements are present, and the percentage (say of sulphides and sulphates) can in many cases be determined; but this is not always required. When the difference between the sum total and the elements found is small, it is reported as "oxygen and loss." When, however, it is considerable, the oxygen may be reported as such; and its amount be either determined directly in the way already described, or calculated from the best determination that can be made of the relative amounts of oxides, sulphides, sulphates, &c., present. Such cases require a careful qualitative analysis to find out that the substance is present; and then the separation of each constituent is made as strictly as possible. These remarks apply especially to ores of the heavy metals. The separation of the constituents is effected with suitable solvents applied in proper order. The soluble sulphates, for example, are extracted with water; the oxides by the dilute acids or alkalies in which they are known to be soluble. The oxygen in the sulphates and oxides thus obtained is estimated by determining the sulphur and metals in the solutions, and calculating the amount of oxygen with which they combine. The metals of the earths and alkalies are almost invariably present as oxides, and are reported as such; except it is known that they are present in some other form, such as fluoride or chloride. Thus, silica, alumina, lime, water, &c., appear in an analysis; even in those cases where "oxygen and loss" is also mentioned. As an example of such a report, take the following analysis of Spanish pyrites:--
Sulphur 49.00 Iron 43.55 Copper 3.20 Arsenic 0.47 Lead 0.93 Zinc 0.35 Lime 0.10 Silica, &c. 0.63 Water 0.70 Oxygen and loss 1.07 ----- 100.00
The following example will illustrate the mode of calculating and reporting. A mineral, occurring as blue crystals soluble in water, and found on testing to be a mixed sulphate of iron and copper, gave on analysis the following results:--
Water 44.51 per cent. Sulphuric oxide 28.82 " Copper 8.44 " Ferrous iron 11.81 " Ferric iron 0.38 " Zinc 0.28 " ----- 94.24
There is here a deficiency of 5.76 per cent. due to oxygen. Nothing else could be found, and it is known that in the sulphates the metals exist as oxides. By multiplying the weight of the copper by 1.252, the weight of copper oxide (CuO) will be ascertained; in this case it equals 10.57 per cent. The ferrous iron multiplied by 1.286 will give the ferrous oxide (FeO); in this case 15.19 per cent. The ferric iron multiplied by 1.428 will give the ferric oxide (Fe_{2}O_{3}); in this case 0.54 per cent. The zinc multiplied by 1.246 will give the zinc oxide (ZnO); in this case it equals 0.35 per cent. The analysis will be reported as--
Water 44.51 Sulphuric oxide 28.82 Copper oxide 10.57 equal to copper 8.44% Ferrous oxide 15.19 Ferric oxide 0.54 Zinc oxide 0.35 ----- 99.98
The following (A) is an analysis of a sample of South American copper ore, which will serve as a further illustration. The analysis showed the presence of 6.89 per cent. of ferrous oxide, and some oxide of copper.
The analysis (B) is that of an ore from the same mine after an imperfect roasting. It will be seen that the carbonates have been converted into sulphates. If the total sulphur simply had been determined, and the sulphate overlooked, the "oxygen and loss" would have been 5.65 per cent., an amount which would obviously require an explanation.
A. B. Water 0.25 0.59 Organic matter 0.54 -- Sulphur 29.50 21.33 Copper 10.92 9.80 {Copper 9.57 {Copper oxide 0.28 Iron 32.09 39.73 {Iron 34.32 {Ferric oxide 7.73 Lead 0.35 0.12 Zinc 0.86 0.69 Cobalt 0.06 0.11 Lime 5.25 7.69 Magnesia 2.33 2.55 Sulphuric oxide 1.00 5.30 Carbon dioxide 8.87 -- "Insoluble silicates" 5.12 8.38 Oxygen and loss 2.86 2.47 ----- Potash 0.15 100.00 Soda 1.09 ----- 100.00
WATER.
Water occurs in minerals in two forms, free and combined. The term "moisture" ought, strictly, to be limited to the first, although, as has already been explained, it is more convenient in assaying to apply the term to all water which is driven off on drying at 100° C. The combined water is really a part of the mineral itself, although it may be driven off at a high temperature, which varies with the base. In some cases a prolonged red heat is required; whilst with crystallised salts it is sometimes given off at the ordinary temperatures. This latter phenomenon, known as efflorescence, is mostly confined to artificial salts.
The determination of the combined water may often be made by simply igniting the substance from which the moisture has been removed. The quantity of water may be determined, either indirectly by the loss, or directly by collecting it in a calcium chloride tube, and weighing. In some cases, in which the loss on ignition does not give simply the proportion of combined water, it can be seen from the analysis to what else the loss is due; and, after a proper deduction, the amount of water can be estimated. For example, 1 gram of crystallised iron sulphate was found to contain on analysis 0.2877 gram of sulphuric oxide; and on igniting another gram, 0.2877 gram of ferric oxide was left. As the salt is known to be made up of ferrous oxide, sulphuric oxide, and combined water, the combined water can be thus calculated: 0.2877 gram of ferric oxide is equal to 0.2589 gram of ferrous oxide,[98] and consequently, the loss on ignition has been diminished by 0.0288 gram, which is the weight of oxygen absorbed by the ferrous oxide during calcining. The loss on ignition was 0.7123 gram, to which must be added 0.0288 gram; hence 0.7411 gram is the weight of the combined sulphuric oxide and water present. Deducting the weight of sulphuric oxide found, 0.2877 gram, there is left for combined water 0.4534 gram. The composition of 1 gram of the dry salt is then:--
Water 0.4534 Sulphuric oxide 0.2877 Ferrous oxide 0.2589 ------ 1.0000
The following is another example:--A sample of malachite lost on ignition 28.47 per cent., leaving a residue which was found on analysis to be made up of oxide of copper (equal to 70.16 per cent. on the mineral), and silica and oxide of iron (equal to 1.37 per cent.). Carbon dioxide and water (but nothing else) was found to be present, and the carbon dioxide amounted to 19.64 per cent.; deducting this from the loss on ignition, we have 8.82 as the percentage of water present. The analysis was then reported as follows:--
Cupric oxide 70.16 equal to 56.0% copper. Silica and ferric oxide 1.37 Carbon dioxide 19.64 Water 8.82 ----- 99.99
~Direct Determination of Combined Water.~--Transfer about 3 grams of the substance to a piece of combustion tube (8 or 10 inches long), attached (as in fig. 63) at one end to a ~U~-tube containing sulphuric acid, and at the other end to a calcium chloride tube. The last is weighed previous to the determination. The tube should be warmed to ensure complete dryness, and must be free from a misty appearance. Aspirate a current of air through the apparatus, heat the mineral by means of a Bunsen burner, cautiously at first, and afterwards to redness (if necessary). The water is driven off and condenses in the calcium chloride tube, which is afterwards cooled and weighed. The increase in weight is due to the water. If the substance gives off acid products on heating, it is previously mixed with some dry oxide of lead or pure calcined magnesia.
EXAMINATION OF WATERS.
The assayer is occasionally called on to test water for the purpose of ascertaining the nature and quantity of the salts contained in it, and whether it is or is not fit for technical and drinking purposes.
In mineral districts the water is generally of exceptional character, being more or less charged, not only with earthy salts, but also frequently with those of the metals. Distilled water is only used by assayers in certain exceptional cases, so that by many it would be classed among the rarer oxides. Water of ordinary purity will do for most purposes, but the nature and quantity of the impurities must be known.
The following determinations are of chief importance:--
~Total Solids at 100° C.~--Where simply the amount is required, take 100 c.c. and evaporate on the water-bath in a weighed dish; then dry in the water-oven, and weigh.
~Total Solids Ignited.~--The above residue is very gently ignited (keeping the heat well below redness), and again weighed. A larger loss than 4 or 5 parts per 100,000 on the water requires an explanation.
~Chlorine.~--Take 100 c.c. of the water in a porcelain dish, add 2 c.c. of a 5 per cent. solution of neutral potassic chromate, and titrate with a neutral standard solution of nitrate of silver, made by dissolving 4.789 grams of crystallised silver nitrate in distilled water, and diluting to 1 litre. The addition of the nitrate of silver is continued until the yellow of the solution assumes a reddish tint. The reaction is very sharp. Each c.c. of nitrate of silver used is equal to 1 part by weight of chlorine in 100,000 of water. At inland places this rarely amounts to more than 1 in 100,000; but near the sea it may amount to 3 or 5. More than this requires explanation, and generally indicates sewage pollution.
~Nitric Pentoxide (N_{2}O_{5}).~--It is more generally reported under the heading, "nitrogen as nitrates." Take 250 c.c. of the water and evaporate to 2 or 3 c.c.; acidulate with a few drops of dilute sulphuric acid, and transfer to a nitrometer (using strong sulphuric acid to wash in the last traces). The sulphuric acid must be added to at least twice the bulk of the liquid. Shake up with mercury. The mercury rapidly flours, and nitric oxide is given off (if any nitrate is present). The volume of the nitric oxide (corrected to normal temperature and pressure), multiplied by 0.25, gives the parts of nitrogen per 100,000; or, multiplied by 0.965, will give the nitric pentoxide in parts per 100,000. In well and spring waters the nitrogen may amount to 0.3 or 0.4 parts per 100,000; or in richly cultivated districts 0.7 or 0.8 parts per 100,000. An excess of nitrates is a suspicious feature, and is generally due to previous contamination.
~Ammonia.~--Take 500 c.c. of the water and place them in a retort connected with a Liebig's condenser. Add a drop or two of a solution of carbonate of soda and distil over 100 c.c.; collect another 50 c.c. separately. Determine the ammonia in the distillate colorimetrically (with Nessler's solution, as described under Ammonia) and compare with a standard solution of ammonic chloride containing 0.0315 gram of ammonic chloride in 1 litre of water. One c.c. contains 0.01 milligram of ammonia. The second distillate will show little, if any, ammonia in ordinary cases. The amounts found in both distillates are added together, and expressed in parts per 100,000.
Waters (other than rain and tank waters) which contain more than 0.003 per 100,000 are suspicious.
~Organic Matter.~--The organic matter cannot be determined directly; but for ordinary purposes it may be measured by the amount of permanganate of potassium which it reduces, or by the amount of ammonia which it evolves on boiling with an alkaline permanganate of potassium solution.
A. Albuminoid Ammonia.--To the residue left after distilling the ammonia add 50 c.c. of a solution made by dissolving 200 grams of potash and 8 grams of potassium permanganate in 1100 c.c. of water, and rapidly boiling till the volume is reduced to 1 litre (this should be kept in a well stoppered bottle, and be occasionally tested to see that it is free from ammonia). Continue the distillation, collecting 50 c.c. at a time, until the distillate is free from ammonia. Three or four fractions are generally sufficient. Determine the ammonia colorimetrically as before. If the total albuminoid ammonia does not exceed 0.005 in 100,000, the water may be regarded as clean as regards organic matter; if it amounts to more than 0.015, it is dirty.
B. Oxygen Consumed.--A standard solution of permanganate of potash is made by dissolving 0.395 gram of the salt in water and diluting to 1 litre. Each c.c. equals 0.1 milligram of available oxygen. The following are also required:--1. A solution of sodium hyposulphite containing 1 gram of the salt (Na_{2}S_{2}O_{3}.5H_{2}O) in 1 litre of water. 2. Dilute sulphuric acid, made by adding one part of the acid to three of water, and titrating with the permanganate solution till a faint pink persists after warming for several hours. 3. Starch paste. 4. Potassium iodide solution.
Take 250 c.c. of the water in a stoppered bottle, add 10 c.c. of sulphuric acid and 10 c.c. of the permanganate, and allow to stand in a warm place for four hours. Then add a few drops of the solution of potassium iodide, and titrate the liberated iodine with "hypo," using starch paste towards the end as an indicator. To standardise the hyposulphite, take 250 c.c. of water and 10 c.c. of sulphuric acid, and a few drops of potassium iodide; then run in 10 c.c. of the "permanganate" solution, and again titrate; about 30 c.c. of the "hypo" will be used. The difference in the two titrations, divided by the last and multiplied by 10, will give the c.c. of permanganate solution used in oxidising the organic matter in the 250 c.c. of water. Each c.c. represents 0.04 parts of oxygen in 100,000.
~Metals.~--These may for the most part be estimated colorimetrically.
~Lead.~--Take 100 c.c. of the water in a Nessler tube, and add 10 c.c. of sulphuretted hydrogen water, and compare the tint, if any, against a standard lead solution, as described under Colorimetric Lead. Report in parts per 100,000.
~Copper.~--Proceed as with the last-mentioned metal; but, if lead is also present, boil down 500 c.c. to about 50 c.c., then add ammonia, filter, and estimate the copper in the blue solution, as described under Colorimetric Copper.
~Iron.~--Take 50 c.c., or a smaller quantity (if necessary), dilute up to the mark with distilled water, and determine with potassium sulphocyanate, as described under Colorimetric Iron.
~Zinc.~--Zinc is the only other metal likely to be present; and, since it cannot be determined colorimetrically, it must be separately estimated during the examination of the "total solids."
~Examination of "Total Solids."~--Evaporate 500 c.c. to dryness with a drop or two of hydrochloric acid. Take up with hydrochloric acid, filter, ignite, and weigh the residue as "silica." To the filtrate add a little ammonic chloride and ammonia, boil and filter, ignite, and weigh the precipitate as "oxide of iron and alumina." Collect the filtrate in a small flask, add a few drops of ammonium sulphide or pass sulphuretted hydrogen, cork the flask, and allow to stand overnight; filter, wash, and determine the zinc gravimetrically as oxide of zinc. If copper or lead were present, they should have been previously removed with sulphuretted hydrogen in the acid solution. To the filtrate add ammonic oxalate and ammonia, boil for some time, allow to stand, filter, wash, ignite, and weigh as "lime." Evaporate the filtrate with nitric acid, and ignite. Take up with a few drops of dilute hydrochloric acid, add baric hydrate in excess, evaporate, and extract with water. The residue contains the magnesia; boil with dilute sulphuric acid, filter, precipitate it with phosphate of soda and ammonia, and weigh as pyrophosphate. The aqueous extract contains the alkalies with the excess of barium. Add sulphuric acid in slight excess, filter, evaporate, and ignite strongly. The residue consists of the sulphates of the alkalies (which are separately determined, as described under Potash).
~Sulphuric Oxide (SO_{3}).~--Take 200 c.c. and boil to a small bulk with a little hydrochloric acid, filter (if necessary), add baric chloride solution in slight excess to the hot solution, filter, ignite, and weigh as baric sulphate.
~Carbon Dioxide (free).~--Carbon dioxide exists in waters in two forms, free and combined. The latter generally occurs as bicarbonate, although on analysis it is more convenient to consider it as carbonate, and to count the excess of carbon dioxide with the free. The method is as follows:--To determine the free carbon dioxide, take 100 c.c. of the water, place them in a flask with 3 c.c. of a strong solution of calcium chloride and 2 c.c. of a solution of ammonic chloride, next add 50 c.c. of lime-water. The strength of the lime-water must be known. Make up to 200 c.c. with distilled water, stop the flask, and allow the precipitate to settle. Take out 100 c.c. of the clear solution with a pipette, and titrate with the standard solution of acid.[99] The number of c.c. required, multiplied by two, and deducted from that required for the 50 c.c. of lime-water, and then multiplied by 0.0045, will give the carbon dioxide present other than as normal carbonates.
~Carbon Dioxide combined~ as normal carbonate.--100 c.c. of the water are tinted with phenacetolin or lacmoid; then heated to near boiling, and titrated with standard acid. The number of c.c. used, multiplied by 0.0045, will give the weight in grams of the combined carbon dioxide.
~Free Acid.~--In some waters (especially those from mining districts) there will be no carbonates. On the contrary, there may be free mineral acid or acid salts. In these cases it is necessary to determine the amount of acid (other than carbon dioxide) present in excess of that required to form normal salts. This is done in the following way:--Make an ammoniacal copper solution by taking 13 grams of copper sulphate (CuSO_{4}.5H_{2}O), dissolving in water, adding solution of ammonia until the precipitate first formed has nearly dissolved, and diluting to 1 litre. Allow to settle, and decant off the clear liquid. The strength of this solution is determined by titrating against 10 or 20 c.c. of the standard solution of sulphuric acid (100 c.c. = 1 gram H_{2}SO_{4}). The finishing point is reached as soon as the solution becomes turbid from precipitated cupric hydrate. At first, as each drop falls into the acid solution, the ammonia and cupric hydrate combine with the free acid to form ammonic and cupric sulphates; but as soon as the free acid is used up, the ammonia in the next drop not only precipitates an equivalent of cupric hydrate from the solution, but also throws down that carried by itself. This method is applicable in the presence of metallic sulphates other than ferric. The standardising and titration should be made under the same conditions. Since sulphuric acid and sulphates are predominant in waters of this kind, it is most convenient to report the acidity of the water as equivalent to so much sulphuric acid.
~Dissolved Oxygen.~--For the gasometric method of analysing for dissolved oxygen, and for the Schützenberger's volumetric method, the student is referred to Sutton's "Volumetric Analysis." The following is an easy method of estimating the free oxygen in a water:--Take 20 c.c. of a stannous chloride solution (about 20 grams of the salt with 10 c.c. of hydrochloric acid to the litre); add 10 c.c. of hydrochloric acid, and titrate in an atmosphere of carbon dioxide with standard permanganate of potassium solution (made by dissolving 1.975 gram of the salt in 1 litre of water: 1 c.c. equals 0.5 milligram of oxygen). A similar titration is made with the addition of 100 c.c. of the water to be tested. Less permanganate will be required in the second titration, according to the amount of oxygen in the water; and the difference, multiplied by 0.5, will give the weight of the oxygen in milligrams. Small quantities of nitrates do not interfere.
In REPORTING the results of the analysis, it is customary to combine the acids and bases found on some such principle as the following:--The sulphuric oxide is calculated as combined with the potash, and reported as potassic sulphate (K_{2}SO_{4}); the balance of the sulphuric oxide is then apportioned to the soda, and reported as sulphate of soda (Na_{2}SO_{4}); if any is still left, it is reported as calcium sulphate (CaSO_{4}), and after that as magnesic sulphate (MgSO_{4}). When the sulphuric oxide has been satisfied, the chlorine is distributed, taking the bases in the same order, then the nitric pentoxide, and lastly the carbon dioxide. But any method for thus combining the bases and acids must be arbitrary and inaccurate. It is extremely improbable that any simple statement can represent the manner in which the bases and acids are distributed whilst in solution; and, since different chemists are not agreed as to any one system, it is better to give up the attempt, and simply state the results of the analysis. This has only one inconvenience. The bases are represented as oxides; and, since some of them are present as chlorides, the sum total of the analysis will be in excess of the actual amount present by the weight of the oxygen equivalent to the chlorine present as chloride. The following is an example of such a statement:--
Parts per 100,000. Total solids, dried at 100° C. 28.73 Chlorine 1.70 Nitrogen as nitrate 0.03 Ammonia 0.001 Albuminoid ammonia 0.004 "Oxygen consumed" in 4 hours 0.01
The solids were made up as under:--
Per 100,000 of the Water. Potash 0.38 Soda 2.01 Magnesia 1.44 Lime 10.55 Ferric oxide 0.01 Silica 0.30 Sulphuric oxide 3.69 Nitrogen pentoxide 0.11 Carbon dioxide 8.38 Chlorine 1.70 Volatile and organic matter 0.66 ----- 29.23 Less oxygen equivalent to chlorine found 0.39 ----- 28.84
For the preparation of distilled water, the apparatus shown in fig. 64 is convenient for laboratory use. It consists of a copper retort heated by a ring gas-burner, and connected with a worm-condenser.
PRACTICAL EXERCISE.
A mineral, on analysis, gave the following results:--Water, 44.94 per cent.; sulphuric oxide, 28.72 per cent.; ferrous iron, 13.92 per cent.; ferric iron, 0.35 per cent.; copper, 6.1 per cent. The mineral was soluble in water, and showed nothing else on testing. How would you report the analysis? Calculate the formula for the salt.
THE HALOGENS.
There is a group of closely allied elements to which the name halogen (salt-producer) has been given. It comprises chlorine, bromine, iodine, and fluorine. These elements combine directly with metals, forming as many series of salts (chlorides, bromides, iodides, and fluorides), corresponding to the respective oxides, but differing in their formulæ by having two atoms of the halogen in the place of one atom of oxygen. For example, ferrous oxide is FeO and ferrous chloride is FeCl_{2}, and, again, ferric oxide is Fe_{2}O_{3}, whilst ferric chloride is Fe_{2}Cl_{6}. These salts differ from the carbonates, nitrates, &c., in containing no oxygen. Consequently, it is incorrect to speak of such compounds as chloride of potash, fluoride of lime, &c., since potash and lime are oxides. It is important to bear this in mind in reporting analyses in which determinations have been made, say, of chlorine, magnesia, and potash, or of fluorine, silica, and alumina. It is necessary in all such cases to deduct from the total an amount of oxygen equivalent to the halogen found, except, of course, where the base has been determined and recorded as metal. Compounds containing oxides and fluorides, &c., do not lend themselves to the method of determining the halogen by difference. For example, topaz, which, according to Dana, has the formula Al_{2}SiO_{4}F_{2}, would yield in the ordinary course of analysis--
Alumina 55.4% Silica 32.6 Fluorine 20.6 ----- 108.6
The oxygen equivalent to 20.6 per cent. fluorine may be found by multiplying the percentage of fluorine by 0.421; it is 8.7 per cent., and must be deducted. The analysis would then be reported thus:--
Alumina 55.4% Silica 32.6 Fluorine 20.6 ----- 108.6 Less oxygen equivalent to fluorine 8.7 ----- 99.9
Take as an illustration the following actual analysis by F.W. Clarke and J.S. Diller:--
Alumina 57.38% Silica 31.92 Fluorine 16.99 Potash 0.15 Soda 1.33 Water 0.20 ------ 107.97 Deduct oxygen equivalent 7.16 ------ 100.81
In calculating the factor for the "oxygen equivalent," divide the weight of one atom of oxygen (16) by the weight of two atoms of the halogen; for example, with chlorine it would be 16/71 or 0.2253; with bromine, 16/160 or 0.1000; with iodine, 16/254 or 0.063; and with fluorine, 16/38 or 0.421.
CHLORINE AND CHLORIDES.
Chlorine occurs in nature chiefly combined with sodium, as halite or rock salt (NaCl). With potassium it forms sylvine (KCl), and, together with magnesium, carnallite (KCl.MgCl_{2}.6H_{2}O). Of the metalliferous minerals containing chlorine, kerargyrite, or horn silver (AgCl), and atacamite, an oxychloride of copper (CuCl_{2}.3Cu(HO)_{2}.) are the most important. Apatite (phosphate of lime) and pyromorphite (phosphate of lead) contain a considerable amount of it. Chlorine is a gas of a greenish colour, possessing a characteristic odour, and moderately soluble in water. It does not occur native, and is generally prepared by the action of an oxidising agent on hydrochloric acid. It combines directly with metals at the ordinary temperature (even with platinum and gold), forming chlorides, which (except in the case of silver) are soluble.
It is important in metallurgy, because of the extensive use of it in extracting gold by "chloridising" processes. It is also used in refining gold.
~Detection.~--Compounds containing the oxides of chlorine are not found in nature, because of the readiness with which they lose oxygen. By reduction they yield a chloride; the form in which chlorine is met with in minerals. In testing, the compound supposed to contain a chloride is boiled with water, or, in some cases, dilute nitric acid. To the clear solution containing nitric acid a few drops of nitrate of silver solution are added. If, on shaking, a white curdy precipitate, soluble in ammonia, separates out, it is sufficiently satisfactory evidence of the presence of chlorides.
~Solution and Separation.~--The chlorides are generally soluble in water, and are got into solution by extracting with warm dilute nitric acid. Or, if insoluble, the substance is fused with carbonate of soda, extracted with water, and the filtrate acidified with nitric acid. For the determination, it is not necessary to obtain the solution of the chloride free from other acids or metals. If tin, antimony, mercury, or platinum is present, it is best to separate by means of sulphuretted hydrogen. The chloride is determined in the solution after removal of the excess of the gas. Where traces of chlorides are being looked for, a blank experiment is made to determine the quantity introduced with the reagents. One hundred c.c. of ordinary water contains from 1 to 3 milligrams of chlorine. On the addition of nitrate of silver to the nitric acid solution, chloride of silver separates out. This is free from other substances, except, perhaps, bromide and iodide.
GRAVIMETRIC DETERMINATION.
Freely mix the solution containing the chloride with dilute nitric acid, filter (if necessary), and treat with nitrate of silver. Heat nearly to boiling, and, when the precipitate has settled, filter, and wash with hot distilled water. Dry, and transfer to a weighed Berlin crucible. Burn the filter-paper separately, and convert any reduced silver into chloride by alternate treatment with drops of nitric and of hydrochloric acid. Add the main portion to this, and heat cautiously till the edges of the mass show signs of fusing (about 260°). Cool in the desiccator and weigh. The substance is chloride of silver (AgCl), and contains 24.73 per cent. of chlorine.
The precipitated chloride is filtered and washed as soon as possible after settling, since on exposure to light it becomes purple, and loses a small amount of chlorine.
VOLUMETRIC METHOD.
There are several volumetric methods; but that based on the precipitation of silver chloride in neutral solution, by means of a standard solution of silver nitrate (using potassium chromate as indicator), is preferred. Silver chromate is a red-coloured salt; and, when silver nitrate is added to a solution containing both chloride and chromate, the development of the red colour marks off sharply the point at which the chloride is used up. Silver chromate is decomposed and consequently decolorised by solution of any chloride. The solution for this method must be neutral, since free acid prevents the formation of the red silver chromate. If not already neutral, it is neutralised by titrating cautiously with a solution of soda. In a neutral solution, other substances (such as phosphates and arsenates) also yield a precipitate with a solution of nitrate of silver; and will count as chloride if they are not removed.
The Standard Solution of Nitrate of Silver is made by dissolving 23.94 grams of the salt (AgNO_{3}) in distilled water, and diluting to 1 litre; 100 c.c. are equal to 0.5 gram of chlorine.
The indicator is made by adding silver nitrate to a strong neutral solution of yellow chromate of potash (K_{2}CrO_{4}), till a permanent red precipitate is formed. The solution is allowed to settle, and the clear liquid decanted into a stoppered bottle labelled "chromate indicator for chlorine."
Standardise the silver nitrate by weighing up 0.5 gram of pure sodium chloride (or potassium chloride). Transfer to a flask and dissolve in distilled water; dilute to 100 c.c. Fill an ordinary burette with the standard silver solution, and (after adjusting) run into the flask a quantity sufficient to throw down the greater part of the chlorine. Add a few drops of the chromate indicator and continue the addition of the silver nitrate until the yellow colour of the solution becomes permanently tinted red, after shaking. This shows that the chlorine is all precipitated, and that the chromate is beginning to come down. The further addition of a couple of drops of the silver solution will cause a marked difference in the tint. Read off the quantity run in, and calculate the standard. One gram of sodium chloride contains 0.6062 gram of chlorine; and 1 gram of potassium chloride contains 0.4754 gram.
For the determination of small quantities of chloride (a few milligrams), the same method is used; but the standard solution is diluted so that each c.c. is equal to 1 milligram of chlorine; and the chromate indicator is added before titrating. The standard solution is made by measuring off 200 c.c. of the solution described above, and diluting with distilled water to 1 litre.
BROMINE AND BROMIDES.
Bromine closely resembles chlorine in the nature of its compounds. It does not occur free in nature, but is occasionally found in combination with silver as bromargyrite (AgBr) and, together with chloride, in embolite. It mainly occurs as alkaline bromides in certain natural waters. Nearly all the bromine of commerce is derived from the mother liquors of salt-works--i.e., the liquors from which the common salt has been crystallised out. Bromine combines directly with the metals, forming a series of salts--the bromides. In ordinary work they are separated with, and (except when specially tested for) counted as, chlorides. They are detected by adding chlorine water to the suspected solution and shaking up with carbon bisulphide. Bromine colours the latter brown.
IODINE AND IODIDES.
Iodine does not occur in nature in the free state; and iodides are rare, iodargyrite or iodide of silver (AgI) being the only one which ranks as a mineral species. Iodates are found associated with Chili saltpetre, which is an important source of the element.
Iodine and Iodides are largely used in the laboratory, and have already been frequently referred to. It is used as an oxidising agent in a similar manner as permanganate and bichromate of potash, especially in the determinations of copper, arsenic, antimony, and manganese.
Iodine is not readily soluble in water; but dissolves easily in a concentrated solution of potassium iodide. Its solutions are strongly coloured; a drop of a dilute solution colours a large volume of water decidedly yellow; on the addition of starch paste, this becomes blue. The delicacy of this reaction is taken advantage of in titrations to determine when free iodine is present. The blue colour may be alternately developed and removed by the addition of iodine (or an oxidising agent) and hyposulphite of soda (or some other reducing agent). In decolorising, the solution changes from blue or black to colourless or pale yellow according to circumstances. Sometimes the solution, instead of remaining colourless, gradually develops a blue which recurs in spite of the further addition of the reducing agent. In these cases the conditions of the assay have been departed from, or (and this is more often the case) there is some substance present capable of liberating iodine.
Iodine forms a series of salts--the iodides--resembling in many respects the chlorides. These can be obtained by direct combination of the metals with iodine.
~Detection.~--Free iodine is best recognised by the violet vapours evolved from the solution on heating, and by the blue or black colour which it strikes on the addition of starch paste. Iodides are detected by boiling with strong solutions of ferric sulphate or chloride. Iodine is liberated, distilled over, and collected. Chlorine also liberates iodine from iodides; and this reaction is frequently made use of in assaying. A process based on this is described under Manganese. All substances which liberate chlorine on boiling with hydrochloric acid (dioxides, bichromates, permanganates, &c.) are determined in a similar way.
~Solution and Separation.~--Most iodides are soluble in water or dilute acids. The separation is effected by distilling the substance with solution of ferric sulphate, and collecting the vapour in a dilute solution of sulphurous acid or arsenite of soda. On the completion of the distillation, the iodine will be in the distillate as iodide; and the gravimetric determination is made on this.
GRAVIMETRIC DETERMINATION.
To the solution containing the iodine, as iodide, and which is free from chlorides (and bromides), add a little dilute nitric acid and nitrate of silver till no further precipitate is produced. Filter off, wash with hot water, and dry. Clean the filter-paper as much as possible, and burn it. Collect the ash in a weighed porcelain crucible, add the main portion, and heat to incipient fusion; cool, and weigh. The substance is silver iodide, and contains 54.03 per cent. of iodine.
VOLUMETRIC METHOD.
This is for the titration of free iodine, and is practically that which is described under Manganese. The substance to be determined is distilled with ferric sulphate, and the iodine is collected in a solution of potassium iodide, in which it readily dissolves. If flaky crystals separate out in the receiver, more potassium iodide crystals are added. When the distillation is finished, the receiver is disconnected, and its contents washed out into a beaker and titrated with "hypo." The standard solution of "hypo" is made by dissolving 19.58 grams of hyposulphite of soda (Na_{2}S_{2}O_{3}.5H_{2}O) in water and diluting to 1 litre; 100 c.c. are equal to 1 gram of iodine. To standardise the solution, weigh up 0.25 gram of pure iodine in a small beaker. Add 2 or 3 crystals of potassium iodide; cover with water; and, when dissolved, dilute to 50 or 100 c.c. Titrate, and calculate the standard.
FLUORINE AND FLUORIDES.
Fluorine is frequently met with as calcium fluoride or fluor-spar (CaF_{2}). It occurs less abundantly as cryolite (Na_{3}AlF_{6}), a fluoride of aluminium and sodium, which is used in glass-making. Certain other rarer fluorides are occasionally met with. Fluorine is also found in apatite, and in some silicates, such as topaz, tourmaline, micas, &c.
Hydrofluoric acid is used for etching glass and opening up silicates. It attacks silica, forming fluoride of silicon (SiF_{4}), which is volatile. Silica is by this means eliminated from other oxides, which, in the presence of sulphuric acid, are fixed. The commercial acid is seldom pure, and generally weak; and the acid itself is dangerously obnoxious. The use of ammonium fluoride (or sodium fluoride) and a mineral acid is more convenient. Determinations of this kind are made in platinum dishes enclosed in lead or copper vessels in a well-ventilated place. Fluor-spar is useful as a flux in dry assaying; it renders slags, which would otherwise be pasty, quite fluid. Fluorides generally are fusible, and impart fusibility to substances with which they form weak compounds. Their fluxing action does not depend on the removal of silicon as fluoride.
~Detection.~--Fluorides in small quantity are easily overlooked unless specially sought for. In larger amounts they are recognised by the property hydrofluoric acid has of etching glass. A watch-glass is warmed, and a layer of wax is melted over the convex side. When cold, some lines are engraved on the waxed surface with any sharp-pointed instrument. The substance to be tested is powdered; and moistened, in a platinum dish, with sulphuric acid. The watch-glass is filled with cold water and supported over the dish. The dish is then carefully warmed, but not sufficiently to melt the wax. After a minute or two, the glass is taken off, and the wax removed. If the substance contained fluorine, the characters will be found permanently etched on the glass. An equally good, but more rapid, test is to mix the powdered substance with some silica, and to heat the mixture in a test tube with sulphuric acid. Silicon fluoride is evolved, and, if a moistened glass rod is held in the tube, it becomes coated with a white deposit of silica, formed by the decomposition of the silicon fluoride by the water. This is also used as a test for silica; but in this case the substance is mixed with a fluoride, and the experiment must obviously be carried out in a platinum vessel.
~Separation and Determination.~--The determination of fluorine is difficult. In the case of fluorides free from silicates (such as fluor-spar), it is determined indirectly by decomposing a weighed portion with sulphuric acid, evaporating, igniting, and weighing the residual sulphate. The increase in weight multiplied by 0.655 gives the weight of fluorine.
In the presence of silica this method does not answer, because of the volatilisation of silicon fluoride. In these cases Wöhler adopted the following plan, which resembles that for the indirect determination of carbon dioxide. Mix the weighed substance in a small flask with powdered silica and sulphuric acid. The mouth of the flask is closed with a cork carrying a tube which is filled, the first half with calcium chloride and the second half with pumice coated with dried copper sulphate. The apparatus is weighed quickly, and then warmed till decomposition is complete. A current of dry air is aspirated for a minute or two; and the apparatus again weighed. The loss in weight gives that of the silicon fluoride (SiF_{4}), which, multiplied by 0.7307, gives the weight of fluorine.
Fresenius uses the same reaction; but collects and weighs the silicon fluoride. The finely powdered and dried substance is mixed with ten or fifteen times its weight of ignited and powdered silica. The mixture is introduced into a small dry flask connected on one side with a series of drying-tubes, and on the other with an empty tube (to condense any sulphuric acid). To this last is joined a drying-tube containing chloride of calcium and anhydrous copper sulphate. This is directly connected with a series of three weighed tubes in which the fluoride of silicon is collected. The last of these is joined to another drying-tube. The first weighed tube contains pumice and cotton wool, moistened with water; the second tube contains soda-lime as well as (in the upper half of the second limb) fused calcium chloride between plugs of wool; the third tube is filled half with soda-lime and half with fused calcium chloride. The distilling-flask containing the substance mixed with silica is charged with 40 or 50 c.c. of sulphuric acid, and placed on the hot plate. Alongside it is placed a similar dry flask containing a thermometer, and the temperature in this is kept at 150° or 160° C. A current of air is sent through the tubes during the operation, which takes from one to three hours for from 0.1 to 1 gram of the substance. A correction is made by deducting 0.001 gram for every hour the dried air has been passed through. The increase in weight of the three tubes gives the weight of the silicon fluoride.
Penfield uses a similar arrangement, but passes his silicon fluoride into an alcoholic solution of potassium chloride. Silica and potassium silico-fluoride are precipitated, and hydrochloric acid is set free.[100] The acid thus liberated is titrated, with a standard solution of alkali, in the alcoholic solution, and from the amount of free acid found the fluorine is calculated. The weight of hydrochloric acid (HCl) found, multiplied by 1.562, gives the weight of the fluorine. With this method of working, fewer ~U~-tubes are required. The exit tube from the flask is bent so as to form a small ~V~, which is kept cool in water; this is directly connected with the ~U~-tube containing the alcoholic solution of potassium chloride. The flask with the assay is heated for about two hours, and a current of dry air is aspirated throughout the determination. Fluoride of silicon is a gas not easily condensed to a liquid: but is immediately decomposed by water or moist air.
FOOTNOTES:
[95] This will require two or three hours to thoroughly complete. It is best to powder the oxide first produced, and recalcine.
[96] No magnetic oxide was formed.
[97] For example:--
CaO + 2HCl = CaCl_{2} + H_{2}O.
PbO + H_{2}SO_{4} = PbSO_{4} + H_{2}O.
MgO + 2HNO_{3} = Mg(NO_{3}){2} + H{2}O.
Al_{2}O_{3} + 6HCl = Al_{2}Cl_{6} + 3H_{2}O.
Fe_{2}O_{3} + 3H_{2}SO_{4} = Fe_{2}(SO_{4}){3} + 3H{2}O.
[98] Fe_{2}O_{3}: 2FeO:: 0.2877: 0.2589.
[99] 100 c.c. contain 1 gram of sulphuric acid.
[100] 3SiF_{4} + 4KCl + 2H_{2}O = 2K_{2}SiF_{6} + SiO_{2} + 4HCl.
XVI. SULPHUR AND SULPHATES
Sulphur occurs native in volcanic districts, and is mined in Sicily, Italy, and California in considerable quantities. Combined with metals (sulphides), it is common in all mineral districts. Iron pyrites (FeS_{2}) is the most abundant source of this element. Sulphates, such as gypsum, are fairly common, but have no value so far as the sulphur in them is concerned. In coal it exists as an impurity, occurring partly as a constituent of organic compounds.
Sulphur, whether free or combined with metals, forms, on burning, sulphurous oxide (SO_{2}), which by the action of oxidising agents and water is converted into sulphuric acid. It forms two oxides, sulphurous (SO_{2}) and sulphuric (SO_{3}), which combine with bases to form sulphites and sulphates. Sulphites are of little importance to the assayer, and are converted into sulphates by the action of nitric acid and other oxidising agents.
The native sulphides, when acted on with hydrochloric acid, give off sulphuretted hydrogen; with nitric acid or aqua regia, sulphates are formed, and more or less sulphur separated.
Sulphur is detected in sulphides by the irritating odour of sulphurous oxide given off on roasting, by the evolution of sulphuretted hydrogen when treated with hydrochloric acid, or by a white precipitate of barium sulphate formed when barium chloride is added to the aqua regia solution.
~Dry Assay.~--There is no method of general application. Free or native sulphur may be volatilised, condensed, and weighed, but pyrites only gives up a portion of its sulphur when heated in a closed vessel, while most sulphides, and all sulphates, give up none at all.
In the determination of sulphur in brimstone, 10 grams of the substance are taken, placed in a small porcelain dish, heated over a Bunsen burner in a well-ventilated place, and ignited. When the sulphur has been completely burnt off, the residue (which consists chiefly of sand) is collected and weighed. In a separate portion the moisture and arsenic are determined; the amounts of these are deducted from the loss in the first experiment. The difference, multiplied by 10, gives the percentage of sulphur.
WET METHODS.
~Solution.~--All sulphates, excepting those of lead, barium, strontium, and lime, are soluble in water or dilute acid. All sulphides, except cinnabar, are converted into sulphates by the action of nitric acid at a gentle heat; or, better, by the action of a mixture of three volumes of nitric acid and one volume of hydrochloric acid. This last attacks cinnabar as well. With most substances it is difficult to convert the whole of the sulphur into sulphuric acid. The sulphur separates out at first as a dark spongy mass, which (on continued treatment) changes to light-coloured flakes. When the solution becomes concentrated and the temperature rises sufficiently, the sulphur fuses into one or more honey-coloured globules which, owing to the small surface they oppose to the acid, are very slowly oxidised. It is not desirable to assist the formation of these globules; therefore, the temperature is kept as low as possible, and strong nitric acid is used. When such globules form, it is best to allow the solution to cool, when the globules will solidify. They can then be filtered off and picked out from the insoluble residue, dried, weighed, ignited, and again weighed, the loss being counted as sulphur. With iron pyrites this difficulty seldom occurs.
Metallic sulphides when fused with an excess of nitre are completely oxidised. If the ore is rich in sulphur, some inert body (such as sodium chloride, or, better, sodium carbonate) is added to dilute the action. With pure sulphur, the action is so energetic as to cause an explosion, so that care should be taken. With burnt ores (incompletely calcined pyrites), there is sufficient oxide of iron present to prevent too rapid action.
These fusions with nitre are best conducted in a platinum dish covered with a piece of platinum foil. The ore is ground with the nitre to ensure complete mixing. The heat need not be excessive, so that a single Bunsen burner placed beneath the dish will suffice; if the bottom of the dish is seen to be red-hot, it is sufficient. On cooling and extracting with water, the sulphur will pass into solution as potassium sulphate, which is then filtered off from the insoluble oxides of iron, copper, &c. The filtrate, after having been treated with a large excess of hydrochloric acid, evaporated to dryness, and re-dissolved in water, is ready for the determination.
Lead sulphate may be dissolved by boiling with ammonium acetate. The insoluble sulphates of barium, strontium, and lime, are decomposed by fusing with 4 or 5 times their weight of "fusion mixture." The alkaline sulphates are then dissolved out with water, and filtered off from the insoluble residue. The filtrate is rendered acid with hydrochloric acid.
~Separation.~--The determination of the sulphuric acid in these solutions by precipitation with barium chloride also serves as a separation; but in hot acid solutions containing copper, and more especially iron salts, the baric sulphate has a strong tendency to carry down amounts of those bodies, varying, no doubt, with the conditions of the precipitation. Boiling hydrochloric acid fails to completely extract them. Moreover, the use of hot concentrated hydrochloric acid causes a loss by dissolving barium sulphate. Nitric acid and nitrates must be decomposed by prolonged boiling and evaporation with hydrochloric acid. The iron may be removed by adding a slight excess of ammonia to the faintly acid solution, filtering off, and washing the precipitated ferric hydrate with hot water. By slightly acidulating the filtrate with hydrochloric acid, it will be rendered ready for the determination.
GRAVIMETRIC METHOD.
This assay is one of those which strikingly shows the necessity of getting the assay solution under proper conditions, in order to obtain satisfactory results. The method has been repeatedly investigated, and the conclusion arrived at, "that it can be correct only by accident." Yet there are many chemists who get good results, and place considerable faith in its accuracy. This can only be due to differences in the manner of working. It is generally understood that nitric acid or nitrates must be absent; and our experience fully confirms this. Precipitations in nitrate solutions are worthless, as the following experiments show. In each experiment the bulk of the solution was 150 c.c. The solutions contained 10 grams of nitre, were freely acid with hydrochloric acid, and were precipitated (while boiling) with slight excess of baric chloride.
Sulphuric acid taken 0.020 gram 0.050 gram 0.100 gram " found 0.019 " 0.047 " 0.098 " " taken 0.500 " 1.004 " 1.000 " " found 0.526 " 1.126 " 1.126 "
All the precipitates were boiled with hydrochloric acid, and thoroughly washed before weighing. The results of some other experiments on this subject are given under "sulphur" in the "examination of commercial copper," page 207.
The solution having been obtained free from nitrates and chlorates (and containing but little free hydrochloric acid), is largely diluted, heated to boiling, and precipitated with a moderate excess of a solution of chloride of barium (8 parts of the crystallized barium chloride are sufficient for 1 of sulphur). It is allowed to settle for half-an-hour, and then decanted through a filter. The precipitate is shaken up with boiling water, rendered slightly acid, filtered, washed, dried, ignited, and weighed. The ignited precipitate, when pure, is white, and is not decomposed at a red heat; it is barium sulphate (BaSO_{4}), and contains 13.73 per cent. of sulphur, or 34.33 per cent. of sulphuric oxide (SO_{3}).
~Determination of Sulphur in Pyrites.~--Weigh up half a gram of the dried and powdered sample, and treat with 10 c.c. of a mixture of 3 volumes of nitric acid and 1 volume of hydrochloric acid, occasionally heating. Evaporate to dryness, treat with 5 c.c. of hydrochloric acid, and again evaporate; take up with 1 c.c. of hydrochloric acid and 100 c.c. of hot water, filter through a small filter, and wash. The residue may contain sulphates of lead, barium, or lime; it must be separately examined, if the total sulphur is wanted. The filtrate is heated, and rendered slightly alkaline with ammonia. Filter off the precipitated ferric hydrate through a quick filter, and wash with hot water. If necessary, evaporate the bulk to about 200 c.c., render faintly acid with hydrochloric acid, and add 20 c.c. of solution of barium chloride; allow to stand for half-an-hour, and decant through a filter. Wash with hot water, dry, ignite, and weigh. Pure pyrites contains 53.33 per cent. of sulphur.
VOLUMETRIC METHOD.
This is based upon the easy conversion of all sulphur compounds into sulphates by fusion with nitre or by oxidation with nitric acid; and on the determination of the sulphate formed by titration in an acetic acid solution with baric chloride.[101] The finishing point is determined by filtering off portions of the assay solution, and testing with sulphuric acid. A slight excess of baric chloride will cause a precipitate.
The process may be divided into--(1) the preparation of the solution, and (2) the titration.
~Preparation of the Solution.~--Weigh up from 1 to 5 grams of the dried and powdered substance, and mix intimately with 4 grams of powdered nitre; clean out the mortar with another gram of nitre, and add this as a cover. Heat in a platinum crucible for fifteen minutes at a low temperature; cool, and extract with water in an evaporating-dish about 9 inches across, and holding 700 or 800 c.c. Add 10 grams of sodium acetate and 10 c.c. of acetic acid, and dilute to half a litre. Boil. The solution is ready for titrating. Substances which lose sulphur on heating (such as pyrites) are thus treated:--Weigh up 1 gram, and evaporate nearly to dryness with 10 c.c. each of nitric and hydrochloric acids. Take up with 10 c.c. of hydrochloric acid, and again boil down to a small bulk; dilute and transfer to a 9-inch evaporating-dish; add 10 grams of sodium acetate and 5 c.c. of acetic acid, dilute to half a litre, and boil. The solution is ready for titrating. Sulphates may be dissolved up in the dish itself with the help of a c.c. or so of hydrochloric acid; sodium acetate and acetic acid are then added; and, after dilution and boiling, the solutions are at once titrated.
The solution before titration must contain no free mineral acid, but 5 or 10 c.c. of acetic acid should be present. It must contain 10 grams of sodium acetate, or sufficient to convert any free mineral acid into its corresponding sodic salt; or, if chlorides, nitrates or sulphates of the metals are present, sufficient to decompose them. If a precipitation occurs, as is the case with ferric salts, &c., the solution is titrated with the precipitate in it.
~The Titration.~--The standard solution of barium chloride is made by dissolving 76.25 grams of the crystallized salt (BaCl_{2}.2H_{2}O) in distilled water, and diluting to 1 litre. 100 c.c. will equal 1 gram of sulphur. As indicator, use dilute sulphuric acid. The strength of the solution may be checked by the titration of 5 grams of ferrous sulphate (oxidized with permanganate of potassium or a few drops of nitric acid), which should require 57.5 c.c. of the barium chloride solution; or any pure sulphate of known composition can be used; anhydrous salts should be preferred.
Fill an ordinary 100 c.c. burette with the solution of barium chloride. The evaporating dish containing the assay solution is placed on a round burner (as shown in fig. 65), and the solution is kept steadily boiling. An ordinary Bunsen-burner flame will cause bumping, and should not be used. Run in the standard solution in quantity known to be insufficient; then withdraw a portion of about 2 c.c., with a pipette, and filter through a fine filter-paper into a test tube. Run in another 0.5 c.c. of the standard solution, and withdraw and filter into a test tube another portion of 2 c.c.; and continue this operation until half-a-dozen or more portions have been drawn off. The test tubes should be arranged in order in a stand resting on a piece of paper, so that each test tube representing 0.5 c.c. of the standard baric chloride may have its value recorded beneath it (fig. 66). Add to each test tube 3 drops of dilute sulphuric acid; that which shows the first appearance of a precipitate marks the point at which the titration is complete. Suppose, for example, that the test tube marked 48.5 c.c. shows no precipitate, while that at 49.0 c.c. shows one, it is evident that the finishing point lies between these readings. With a little practice, one can judge from the appearance of the precipitate in the 49 c.c. tube, whether 1/4 c.c. should be deducted or not.
It is better to add dilute sulphuric acid, and to watch for the appearance of a precipitate in the test tube, than to add baric chloride and to look for its non-appearance; besides, baric chloride is much less likely to be present in a test tube as impurity than sulphates are. In this way the chance of error from what are termed "accidental causes" is diminished.
The following experiments show the effect of variation in the conditions of titration:--
Make a standard solution of sulphuric acid by diluting 43.65 grams of sulphuric acid (sp. g. 1.6165) to 1 litre: 100 c.c. will contain 1 gram of sulphur. An equivalent solution may be made by dissolving 100.62 grams of sodium sulphate crystals (Na_{2}SO_{4}.10H_{2}O), or 86.88 grams of ferrous sulphate (FeSO_{4}.7H_{2}O), in water (oxidising the latter), and diluting to 1 litre.
The order in which these experiments are given is that in which they were made in an investigation into the conditions under which the titration could most accurately be effected.
~Effect of Hydrochloric and Nitric Acids.~--The titrations were performed in the manner already described, but sodic acetate and acetic acid were absent. Twenty c.c. of the standard solution of sulphuric acid were used.
Hydrochloric acid present 0.0 c.c. 1.0 c.c. 2.0 c.c. 5.0 c.c. "Baric chloride" required 20.0 " 20.0 " 19.7 " 12.5 "
Nitric acid present 0.0 c.c. 1.0 c.c. 2.0 c.c. 5.0 c.c. "Baric chloride" required 20.0 " 19.5 " 18.0 " 10.0 "
These show clearly the interference of free mineral acids, although very dilute hydrochloric acid (1 c.c. in 500 of water) has no effect.
~Effect of Acetic and Citric Acids.~--A similar series of experiments with these acids gave the following results:--
Acetic acid present 0.0 c.c. 5.0 c.c. 50.0 c.c. 100.0 c.c. "Baric chloride" required 20.0 " 20.0 " 20.0 " 20.0 "
Citric acid present 0 gram 1 gram 5 grams "Baric chloride" required 20.0 c.c. 20.0 c.c. 20.0 c.c.
These acids do not interfere.
~Effect of Sodic Acetate and Acetic Acid.~--In each of these experiments 5 c.c. of acetic acid was present.
Sodium acetate added 0 gram 1 gram 10 grams 50 grams "Baric chloride" required 20.0 c.c. 20.0 c.c. 20.0 c.c. 20.0 c.c.
As sodic acetate and acetic acid did not interfere, it became desirable to make some experiments on the finishing point. The first object sought for was the smallest amount of the standard baric chloride in 500 c.c. of water, required to give an indication when tested in the manner already described.
Baric Chloride Conditions of Assay Solution. required.
Water only 0.05 c.c. With 10 grams of sodium acetate and 5 c.c. of acetic acid 0.05 " The same with 5 grams of nitre 0.10 " Like the last, but with 5 grams of salt instead of nitre 0.10 "
These show that as small an amount of baric chloride solution as is equal to only 0.000002 gram of sulphur in the 2 c.c. of solution tested yields a decided precipitate on the addition of 3 drops of sulphuric acid.
To determine whether the same finishing point is obtained on testing the filtered portions in the test tubes with baric chloride as is obtained on testing with sulphuric acid, a titration was made with 20 c.c. of standard solution of sulphuric acid, together with the usual quantities of sodic acetate and acetic acid; and two lots of 2 c.c. each were filtered into two sets of test tubes after each addition of the standard baric chloride. To one series 3 drops of baric chloride solution were added, and to the other 3 drops of sulphuric acid. The results were--
With Dilute With Baric "Baric Chloride" added. Sulphuric Acid. Chloride Solution.
19.5 c.c. Clear Cloudy 19.75 " Clear Cloudy 20.0 " Finished Finished 20.25 " Cloudy Clear 20.5 " Cloudy Clear
The two methods of testing give the same result. But this balance is disturbed in the presence of much nitre, the indications with baric chloride being disturbed by an opalescence for some c.c. beyond the finishing point. In solutions containing free hydrochloric or nitric acid, a precipitate is obtained with either baric chloride or sulphuric acid.
~Effect of Varying Sulphur.~--In these and the subsequent experiments the titrations were performed in the presence of 10 grams of sodic acetate and 10 c.c. of acetic acid in the manner already described.
Standard sulphuric acid used 5.0 c.c. 10.0 c.c. 20.0 c.c. 50.0 c.c. 100.0 c.c.
"Baric chloride" required 5.0 " 10.0 " 20.0 " 50.0 " 100.0 "
~Effect of Varying Temperature.~--With 5 c.c. of standard sulphuric acid titrated at 15° C., 5 c.c. of baric chloride were required; but with larger quantities the results were altogether unsatisfactory when titrated cold.
~Effect of Varying Bulk.~--
Bulk 100.0 c.c. 200.0 c.c. 500.0 c.c. 1000.0 c.c. "Baric chloride" required 20.0 " 20.0 " 20.0 " 20.5 "
Considerable variation in bulk has no effect, but 500 c.c. is the most convenient volume to work with. It is well to occasionally replace the water boiled off during titration.
~Effect of Foreign Salts.~--In all these experiments 20 c.c. of "sulphuric acid" were used, and the titration was performed in the ordinary way.
Sodic chloride added 0 gram 5 grams 10 grams "Baric chloride" required 20.0 c.c. 20.0 c.c. 19.7 c.c.
Ammonic chloride added 0 gram 5 grams 10 grams "Baric chloride" required 20.0 c.c. 20.0 c.c. 19.5 c.c.
Calcic chloride added 0 gram 1 gram 2 grams 5 grams "Baric chloride" required 20.0 c.c. 20.0 c.c. 19.2 c.c. 19.0 c.c.
Zinc chloride added 0 gram 1 gram 3 grams 5 grams "Baric chloride" required 20.0 c.c. 20.0 c.c. 20.0 c.c. 20.0 c.c.
Ferrous chloride added 0 gram 1 gram 3 grams 5 grams "Baric chloride" required 20.0 c.c. 19.7 c.c. 19.5 c.c. 19.0 c.c.
Ferric chloride added 0 gram 1 gram 3 grams 5 grams "Baric chloride" required 20.0 c.c. 20.0 c.c. 20.0 c.c. 20.0 c.c.
Copper chloride added 0 gram 1 gram 3 grams 5 grams "Baric chloride" required 20.0 c.c. 20.0 c.c. 20.0 c.c. 20.0 c.c.
Potassic Nitrate added 0 gram 1 gram 5 grams 10 grams "Baric chloride" required 20.0 c.c. 20.0 c.c. 20.0 c.c. 19.0 c.c.
Potassic Nitrite added 0 gram 1 gram 5 grams "Baric chloride" required 20.0 c.c. 20.0 c.c. 20.0 c.c.
Sodic phosphate added 0 gram 1 gram "Baric chloride" required 20.0 c.c. 22.5 c.c.
Sodic arsenate added 0 gram 1 gram "Baric chloride" required 20.0 c.c. 20.5 c.c.
In the absence of ferric salts, phosphates and arsenates count as sulphur.
In two series of experiments for determining the effect of varying amounts of sulphur in the form of ferrous sulphate, we obtained the following results:--In the first series the assay solution was prepared in the manner we have described for Pyrites; and in the second series, by fusion with nitre.
Sulphur added 0.050 gram 0.100 gram 0.200 gram "Baric chloride" required (1) 5.0 c.c. 10.0 c.c. 20.0 c.c. " " (2) 4.7 " 10.0 " 20.0 "
Sulphur added 0.500 gram 1.000 gram "Baric chloride" required (1) 50.0 c.c. 100.0 c.c. " " (2) 50.0 " 100.0 "
More than 5 grams of nitre must not be used in an assay; and, since the requisite amount of nitre considerably exceeds that sufficient to oxidise the sulphur, not more than 0.5 gram of unoxidised sulphur should be present in the portion of the sample weighed up for determination. When the amount of sulphur present is not known within reasonable limits, the test portions may be tried with a drop of baric chloride solution instead of sulphuric acid, so that the diminishing quantity of precipitate may give warning of an approach to the finishing point.
~Determination of Sulphur in Blende.~--Weigh up 1 gram of dried and powdered blende, and mix and fuse with 5 grams of nitre in the manner described. Place the dish and its contents in the titrating-dish, extract with water, add 10 grams of sodium acetate and 10 c.c. of acetic acid, remove and wash the platinum-dish, and dilute to 500 c.c.; boil and titrate. In the example, duplicate determinations required (a) 32.0 c.c., (b) 32.25 c.c., giving an average of 32.1 per cent. of sulphur.
~Determination of Sulphur in Chalcopyrite~ (Yellow Copper Ore).--Take 1 gram of the finely-powdered sample, and 5 grams of nitre. Sprinkle a little of the nitre in a small Wedgwood mortar, place the ore on it, and cover with 2 or 3 grams more of the nitre. Rub up together, and transfer to a small porcelain dish; clean out the mortar with the rest of the nitre, and add to the contents of the dish. Cover with a piece of platinum foil, and heat gently with a Bunsen burner till the nitre melts and the stuff shows signs of deflagrating; remove the heat, and allow the action to go on by itself for a minute or so, then heat over the Bunsen burner for 10 minutes. Cool; transfer the whole to the titrating-dish; boil with 500 c.c. of water; remove the small dish and foil; add sodic acetate and acetic acid, and titrate.
For example, 1 gram required 34.5 c.c. of "barium chloride" (standard = 1.005 gram S), which is equivalent to 34.7 per cent. sulphur. The theoretical percentage is 34.8.
~Determination of Sulphur in Chalcocite~ (Grey Copper Ore).--Proceed as in the last experiment but, since the action with nitre is more moderate, no special precautions need be taken on heating. A platinum dish may be used.
An example which was heated for 30 minutes required 20.5 c.c. of the barium chloride solution. This is equivalent to 20.6 per cent. of sulphur. The theoretical yield is 20.2 per cent.
~Determination of Sulphur in Pyrites.~--Take 1 gram of the finely-powdered sample, cover with 10 c.c. of nitric acid, and, when action has ceased, evaporate to a small bulk. Add 3 or 4 c.c. of hydrochloric acid, and again evaporate to a paste. Take up with 1 or 2 c.c. of dilute hydrochloric acid, dilute with water, transfer to a titrating-dish, add 10 grams of sodic acetate and 5 c.c. of acetic acid, and dilute with water to 500 c.c. Boil and titrate.
An example with 1 gram of a pure crystallized pyrites required 52.7 c.c. of the barium chloride solution, which is equivalent to 53.0 per cent. of sulphur. Theory requires 53.3 per cent. of sulphur.
~Determination of Sulphur in Mispickel.~--Take 1 gram of the powdered ore and evaporate with 10 c.c. of nitric acid, and take up with 3 or 4 c.c. of hydrochloric acid. If any globules of sulphur remain, again evaporate with nitric acid. Dilute, and transfer to the titrating-dish. Add 10 grams of sodic acetate, dilute with water, boil, and titrate. The mispickel carries (according to theory) exactly sufficient iron to precipitate the arsenic as ferric arsenate in an acetic acid solution, so no more iron need be added. The ferric arsenate will separate out as a yellowish-white flocculent precipitate.
An example required, in duplicate experiment, 18.5 c.c. and 18.7 c.c. of barium chloride, equivalent to 18.7 per cent. of sulphur. The formula, FeS_{2}.FeAs_{2}, requires 19.6 percent., but the sulphur generally varies considerably from this amount.
~Determination of Sulphur in Burnt Ores.~--Take 5 grams of the dried and powdered ore, and rub up with 4 grams of nitre; transfer to the platinum-dish; clean out the mortar with another gram of nitre, and add this as a cover. Heat, and extract with water as before; add the sodium acetate and acetic acid; and titrate. Burnt ores carry from 2.5 to 5 per cent. of sulphur. A series of four determinations gave:--
"Baric Chloride" Required. Percentage of Sulphur. Gravimetric Results. 12.6 c.c. 2.52 % 2.45 % 29.9 " 5.98 " 5.84 " 18.1 " 3.62 " 3.53 " 22.0 " 4.40 " 4.43 "
For ores carrying less than 1 per cent. of sulphur, take 10 grams for the assay.
~Determination of Sulphuric Oxide (SO_{3}) in Sulphates.~--When the sulphur exists in the sample received by the assayer in an oxidised state as sulphate, it is usual to report it in terms of sulphuric oxide (SO_{3}). In this case, the metal must also be reported as oxide. For example, an analysis of copper sulphate would be thus reported:--
Oxide of copper (CuO) 31.8 % Sulphuric oxide (SO_{3}) 32.1 " Water 36.1 " ----- 100.0
The percentage of sulphur multiplied by 2.5 gives the percentage of sulphuric oxide. Thus a sample of copper sulphate containing 12.85 per cent. of sulphur will contain 12.85 × 2.5 or 32.12 per cent. of sulphuric oxide.
In minerals and metallurgical products, it is common to find the sulphur in both conditions--i.e., as sulphate and sulphide. Generally in these the percentage of sulphur only is wanted; but this will depend entirely on commercial requirements, and not on the fancy of the assayer. Soluble sulphates are determined separately by extracting with small quantities of cold water, so as to avoid the separation of basic sulphates, or, if the sulphides present are not at the same time attacked, by dilute hydrochloric acid. Lead sulphate may be extracted by boiling with ammonic acetate; whilst barium, strontium, and, perhaps, calcium sulphate, will be mainly found in the residue insoluble in acids.
Weigh up from 2 to 5 grams of the material according to the amount of sulphur judged to be present, and dissolve them in the titrating-dish with 1 c.c. of hydrochloric acid and 50 c.c. of water. Add 10 grams of sodic acetate, and 10 c.c. of acetic acid; dilute, boil, and titrate. In the case of ferric salts, half the quantity of acetic acid will be better, as then the ferric iron will be precipitated, and a colourless solution will be left, in which the end reaction is more readily distinguished.
Determined in this way, 5 gram samples of the following salts gave the results indicated below:--
"Barium Chloride" Salt. Required. Sulphuric Oxide. Copper sulphate 64.25 c.c. 32.12 % Magnesium sulphate 65.25 " 32.62 " Zinc sulphate 56.25 " 28.12 " Ferrous sulphate 58.25 " 29.12 " Sodium sulphate 51.25 " 25.60 "
~Determination of Sulphuric Oxide in Barytes~ (Heavy spar).--Fuse 2 grams of the powdered mineral with 5 grams of "fusion mixture" for five minutes; and, when cold, extract with water. Filter, acidulate the filtrate with an excess of 10 c.c. of acetic acid, dilute, boil, and titrate. For example, a transparent crystallised sample required 27.0 c.c. of barium chloride, which is equivalent to 13.6 per cent. of sulphur, or 34.0 per cent. of sulphuric oxide. Theory requires 34.3 per cent. of the latter. Since both carbonate of soda and potash are liable to contain sulphates, a blank determination should be made on 5 grams of the "fusion mixture," and the amount found be deducted from that got in the assay.
PRACTICAL EXERCISES.
1. The price of sulphur in an ore being 4-1/2d. per unit in the northern markets, what would be the price of a ton of ore containing 49 per cent. of sulphur? What would be the effect on the price of an error of 0.25 per cent. in the assay?
2. Pyrites carries 50 per cent. of sulphur, and on calcining yields 70 per cent. of its weight of burnt ore. Supposing the burnt ore carries 3.5 per cent. of sulphur, what proportion of the sulphur will have been removed in the calcining?
3. How would blende compare with pyrites as a source of sulphur for sulphuric acid making?
4. How would you determine the percentage of sulphuric oxide in a sample of gypsum? What is sulphuric oxide, and what relation does it bear to sulphur?
5. A mineral contains 20.7 per cent. of water, 32.4 per cent. of lime, and 18.6 per cent. of sulphur. What is its probable composition? What experiment would you try to determine the accuracy of your conclusion?
SELENIUM
occurs in nature combined with copper, mercury, and lead, in certain rare minerals. In small quantities it is found in many ores. It is detected in solution by the red precipitate produced on boiling the acid solution with sodium sulphite. This reaction is used for its determination.
~Solution.~--The solution is effected by boiling with nitric acid or aqua regia, or by fusing with nitre. To separate the selenium, the solution is evaporated with an excess of hydrochloric acid and a little sodium or potassium chloride. This destroys any nitric acid that may be present, and reduces selenic acid (H_{2}SeO_{4}) to selenious (H_{2}SeO_{3}). The solution is diluted with water, and treated with a solution of sulphite of soda. It is warmed, and at last boiled. The selenium separates as a red precipitate, which (on boiling) becomes denser and black. It is collected on a weighed filter, washed with hot water, dried at 100° C., and weighed as pure selenium.
Selenium can be precipitated with sulphuretted hydrogen as a sulphide, which is readily soluble in ammonium sulphide. This sulphide may be oxidised with hydrochloric acid and chlorate of potash; and the selenium separated in the manner described.
TELLURIUM.
Tellurium occurs in nature, native, and in combination with gold, silver, bismuth and lead. It is sometimes met with in assaying gold ores. It may be detected by the purple colour it imparts to strong sulphuric acid when dissolved in the cold, and by the black precipitate of metallic tellurium which its solutions yield on treatment with a reducing agent. Telluric acid is reduced to tellurous (with evolution of chlorine) on boiling with hydrochloric acid.
~Solution~ is effected by boiling with aqua regia, or by fusing with nitre and sodium carbonate.
~Separation.~--Tellurium closely resembles selenium in its reactions. It is separated and determined in the same way. Like it, it forms a sulphide soluble in ammonium sulphide. It is distinguished from selenium by the insolubility, in a solution of cyanide of potassium, of the metal precipitated by sodium sulphite; whereas selenium dissolves, forming a soluble potassic seleno-cyanide.[102]
For the determination, solution is effected by fusing with nitre and sodium carbonate, dissolving out the tellurate of potash with water, and boiling with hydrochloric acid. Tellurous compounds are formed, with evolution of chlorine; and the solution, on treating with a reducing agent (such as sulphurous acid or stannous chloride), yields metallic tellurium; which is washed, dried at 100° C., and weighed.
FOOTNOTES:
[101] BaCl_{2} + Na_{2}SO_{4} = BaSO_{4} + 2NaCl.
[102] Se + KCy = KCySe.
XVII. ARSENIC, PHOSPHORUS, NITROGEN
ARSENIC.
The chief source of the arsenic of commerce is arsenical pyrites, or mispickel, which contains about 45 per cent. of arsenic (As). Arsenic also occurs as a constituent of several comparatively rare minerals; and, as an impurity, it is very widely distributed. White arsenic is an oxide of arsenic, and is obtained by roasting arsenical ores, and refining the material (crude arsenic), which condenses in the flues. Arsenic itself is volatile, and many of its compounds have the same property. It forms two well-defined series of salts, corresponding to the oxides: arsenious oxide (As_{2}O_{3}), and arsenic oxide (As_{2}O_{5}). These combine with bases to form arsenites and arsenates respectively. Boiling with nitric acid converts the lower into the higher oxide; and powerful reducing-agents, such as cuprous chloride, have the opposite effect.
Arsenic may be detected by dissolving the substance in hydrochloric acid, or in aqua regia (avoiding an excess of nitric acid), and adding a little of this solution to the contents of a small flask in which hydrogen is being made by the action of zinc and hydrochloric acid. The ignited jet of hydrogen assumes a blue colour if arsenic is present, and a cold porcelain dish held in the flame (fig. 67) becomes coated with a dark deposit of metallic arsenic. Antimony produces a similar effect, but is distinguished by the insolubility of its deposit in a cold solution of bleaching-powder.
Arsenites are distinguished by the volatility of the chloride; by decolorising a solution of permanganate of potassium, and by immediately giving a yellow precipitate with sulphuretted hydrogen. Arsenates are distinguished (after converting into soda salts by boiling with carbonate of soda and neutralising) by giving with nitrate of silver a red precipitate, and with "magnesia mixture" a white crystalline one.
~Dry Assay.~--There is no dry assay which is trustworthy. The following method is sometimes used to find the proportion of arsenious oxide in "crude arsenic":--Weigh up 5 grams of the dried sample, and place them in a clean dry test-tube about 6 inches long. Tie a small filter-paper over the mouth of the tube, so as to prevent air-currents. Heat the tube cautiously so as to sublime off the white arsenic into the upper part of the tube. Cut off the bottom of the test-tube by wetting whilst hot. Scrape out the arsenic and weigh it. The weight gives an approximate idea of the quantity, and the colour of the quality, of the white arsenic obtainable from the sample. Some workers (sellers) weigh the residue, and determine the white arsenic by difference. In determining the percentage of moisture in these samples, the substance is dried on a water-bath or in a water-oven.
WET METHODS.
~Solution.~--Where, as in crude arsenic, the substance is arsenious oxide (As_{2}O_{3}) mixed with impurities, the arsenic is best got into solution by warming with caustic soda, and neutralising the excess with hydrochloric acid; it will be present as sodium arsenite. Metals and alloys are acted on by means of nitric acid; or the arsenic may be at the same time dissolved and separated by distilling with a strongly-acid solution of ferric chloride, in the way described under Volumetric Methods.
With minerals, mattes, &c., solution is thus effected:--The finely-powdered substance is mixed (in a large platinum or porcelain crucible) with from six to ten times its weight of a mixture of equal parts of carbonate of soda and nitre. The mass is then heated gradually to fusion, and kept for a few minutes in that state. When cold, it is extracted with warm water, and filtered from the insoluble residue. The solution, acidified with nitric acid and boiled, contains the arsenic as sodium arsenate. With mispickel, and those substances which easily give off arsenic on heating, the substance is first treated with nitric acid, evaporated to dryness, and then the residue is treated in the way just described.
When the arsenic is present as arsenite or arsenide, distillation with an acid solution of ferric chloride will give the whole of the arsenic in the distillate free from any metal except, perhaps, tin as stannic chloride. With arsenates, dissolve the substance in acid and then add an excess of soda. Pass sulphuretted hydrogen into the solution; warm, and filter. Acidulate the filtrate, and pass sulphuretted hydrogen. Decant off the liquid through a filter, and digest the precipitate with ammonic carbonate; filter, and re-precipitate with hydrochloric acid and sulphuretted hydrogen. Allow to stand in a warm place, and filter off the yellow sulphide of arsenic. Wash it into a beaker, clean the filter-paper (if necessary) with a drop or two of dilute ammonia; evaporate with 10 c.c. of dilute nitric acid to a small bulk; dilute; and filter off the globules of sulphur. The filtrate contains the arsenic as arsenic acid.
GRAVIMETRIC METHOD.
Having got the arsenic into solution as arsenic acid, and in a volume not much exceeding 50 c.c., add about 20 c.c. of dilute ammonia and 20 c.c. of "magnesia mixture." Stir with a glass rod, and allow to settle overnight. Filter, and wash with dilute ammonia, avoiding the use of large quantities of wash water. Dry, transfer the precipitate to a Berlin crucible, and clean the filter-paper thoroughly. Burn this paper carefully and completely; and add the ash to the contents of the crucible, together with 4 or 5 drops of nitric acid. Evaporate with a Bunsen burner, and slowly ignite, finishing off with the blow-pipe or muffle. Cool, and weigh. The ignited precipitate is pyrarsenate of magnesia (Mg_{2}As_{2}O_{7}), and contains 48.4 per cent. of arsenic (As).
Instead of igniting the precipitate with nitric acid, it may be collected on a weighed filter-paper, dried at 100° C., and weighed as ammonic-magnesic arsenate (2AmMgAsO_{4}.H_{2}O), which contains 39.5 per cent. of arsenic. The results in this case are likely to be a little higher. The drying is very tedious, and is likely to leave behind more water than is allowed for in the formula. In a series of determinations in which the arsenic was weighed in both forms, the results were:--
Ammonic-magnesic Arsenic Magnesium Pyrarsenate Arsenic Arsenate in grams. in grams. in grams. in grams. 0.0080 0.0032 0.0065 0.0031 0.0400 0.0158 0.0330 0.0160 0.0799 0.0316 0.0633 0.0306 0.1600 0.0632 0.1287 0.0623 0.4000 0.1580 0.3205 0.1551 0.7990 0.3156 0.6435 0.3114
VOLUMETRIC METHODS.
There are two methods: one for determining the arsenic in the lower, and the other in the higher state of oxidation. In the first-mentioned method this is done by titrating with a standard solution of iodine; and in the latter with a solution of uranium acetate. Where the arsenic already exists as arsenious oxide, or where it is most conveniently separated by distillation as arsenious chloride, the iodine method should be used; but when the arsenic is separated as ammonic-magnesic arsenate or as sulphide, the uranium acetate titration should be adopted.
IODINE PROCESS.
This is based on the fact that sodium arsenite in a solution containing an excess of bicarbonate of soda is indirectly oxidised by iodine to sodium arsenate,[103] and that an excess of iodine may be recognised by the blue colour it strikes with starch. The process is divided into two parts--(1) the preparation of the solution, and (2) the titration.
~Preparation of the Solution.~--For substances like crude arsenic, in which the arsenic is present as arsenious oxide, the method is as follows:--Take a portion which shall contain from 0.25 to 0.5 gram of the oxide, place in a beaker, and cover with 10 c.c. of sodic hydrate solution; warm till dissolved, put a small piece of litmus paper in the solution, and render acid with dilute hydrochloric acid. Add 2 grams of bicarbonate of soda in solution, filter (if necessary), and dilute to 100 c.c. The solution is now ready for titrating.
Where the arsenic has to be separated as arsenious chloride, the process is as follows:[104]--Weigh up 1 gram of the finely-powdered ore (metals should be hammered out into a thin foil or be used as filings), and place in a 16-ounce flask provided with a well-fitting cork, and connected with a ~U~-tube, as shown in the drawing (fig. 68). The ~U~-tube should contain 2 or 3 c.c. of water, and is cooled by being placed in a jar or large beaker of cold water. The water used for cooling should be renewed for each assay.
Pour on the assay in the flask 50 c.c. of a "ferric chloride mixture," made by dissolving 600 grams of calcium chloride and 300 grams of ferric chloride in 600 c.c. of hydrochloric acid, and making up to 1 litre with water.
Firmly cork up the apparatus, and boil over a small Bunsen-burner flame for fifteen or twenty minutes, but avoid evaporating to dryness. Disconnect the flask, and pour away its contents at once to prevent breakage of the flask by their solidification. The arsenic will be condensed in the ~U~-tube, together with the greater part of the hydrochloric acid; transfer the distillate to a beaker washing out the tube two or three times with water; add a small piece of litmus paper; neutralise with ammonia; render faintly acid with dilute hydrochloric acid; add 2 grams of bicarbonate of soda in solution; and dilute to 250 c.c. The solution is now ready for titrating.
The arsenic comes over in the early part of the distillation, as will be seen from the following experiment, made on 1 gram of copper precipitate; in which experiment the distillate was collected in separate portions at equal intervals, and the arsenic in each portion determined:--
Time Iodine Equivalent to Arsenic Distilling. Required. in the Distillate.
5 minutes 12.0 c.c. 0.0450 gram 5 " 0.17 " 0.0005 " 5 " 0.0 " 5 " 0.0 " To dryness 0.0 "
The volume of each distillate was about 5 c.c.
In this operation the metals are converted into chlorides by the action of ferric chloride, which gives up a part of its chlorine, and becomes reduced to the ferrous salt. The calcium chloride does not enter into the chemical reaction, but raises the temperature at which the solution boils, and is essential for the completion of the distillation.[105] Two experiments with material containing 3.48 per cent. of arsenic gave--(1) with ferric chloride alone, 2.74 per cent.; and (2) with the addition of calcium chloride, 3.48 per cent.
It is always necessary to make a blank determination with 1 gram of electrotype copper, to find out the amount of arsenic in the ferric chloride mixture.[106] Unfortunately, a correction is always required. This amounts to about 0.15 per cent. of arsenic on each assay, even when the mixture has been purified; and this constitutes the weakness of the method, since, in some cases, the correction is as much as, or even greater than, the percentage to be determined.
The acid distillate containing the arsenious chloride may be left for an hour or so without much fear of oxidation; but it is safer to neutralise and then to add the bicarbonate of soda, as the following experiments show. Several portions of a solution, each having a bulk of 100 c.c., were exposed for varying lengths of time, and the arsenic in each determined.
_____________________________________________________________________ | | | | | | Acid Solutions. | Neutralised Solutions. | | Time Exposed. | "Iodine" Arsenic Found. |"Iodine" Arsenic Found. | | | Required. | Required | |_______________|___________________________|_________________________| | | | | | -- | 18.2 c.c. = 0.0136 gram | 18.1 c.c. = 0.0136 gram | | 1 hour | 18.2 " = 0.0136 " | 18.2 " = 0.0136 " | | 2 hours | 17.7 " = 0.0133 " | 18.0 " = 0.0135 " | | 4 " | 17.5 " = 0.0131 " | 18.4 " = 0.0138 " | | 5 " | 17.0 " = 0.0127 " | 18.3 " = 0.0137 " | |_______________|___________________________|_________________________|
~The Titration.~--Make a standard solution of iodine by weighing up in a beaker 16.933 grams of iodine and 30 grams of potassium iodide in crystals; add a few c.c. of water, and, when dissolved, dilute to 1 litre: 100 c.c. will equal 0.500 gram of arsenic.
A solution of starch similar to that used in the iodide-copper assay will be required. Use 2 c.c. for each assay. Variations in the quantity of starch used do not interfere; but the solution must be freshly prepared, as after seven or eight days it becomes useless.
To standardise the iodine solution, weigh up 0.3 gram of white arsenic; dissolve in caustic soda; neutralise; after acidulating, add 2 grams of bicarbonate of soda and 2 c.c. of the starch solution, and dilute to 200 c.c. with cold water. Fill a burette having a glass stop-cock with the iodine solution, and run it into the solution of arsenic, rapidly at first, and then more cautiously, till a final drop produces a blue colour throughout the solution. Calculate the standard in the usual way. White arsenic contains 75.76 per cent. of arsenic.
The following experiments show the effect of variation in the conditions of the titration:--
Make a solution of arsenic by dissolving 6.60 grams of white arsenic in 100 c.c. of sodic hydrate solution; render slightly acid with hydrochloric acid; add 10 grains of bicarbonate of soda, and dilute to 1 litre: 100 c.c. will contain 0.50 gram of arsenic.
~Effect of Varying Temperature.~--The reaction goes on very quickly in the cold, and, since there is no occasion for heating, all titrations should therefore be carried out cold.
~Effect of Varying Bulk.~--In these experiments, 20 c.c. of arsenic solution were taken, 2 grams of bicarbonate of soda and 2 c.c. of starch solution added, and water supplied to the required bulk. The results were:--
Bulk 50.0 c.c. 100.0 c.c. 250.0 c.c. 500.0 c.c. "Iodine" required 20.0 " 20.0 " 20.0 " 20.0 "
Considerable variation in bulk does not interfere.
~Effect of Varying Bicarbonate of Soda.~--This salt must be present in each titration in considerable excess, to prevent the interference of free acid. The bicarbonate must be dissolved without heating, as neutral carbonates should be avoided.
Bicarbonate added 1 gram 2 grams 5 grams 10 grams "Iodine" required 20.1 c.c. 20.0 c.c. 20.1 c.c. 20.0 c.c.
These results show that large variation in the quantity of bicarbonate has no effect.
~Effect of Free Acid.~--In these experiments, the arsenic taken, the starch, and the bulk were as before, but no bicarbonate was added. In one case the solution was rendered acid with 5 c.c. of acetic acid, and in the other with 5 c.c. of hydrochloric acid; in both cases the interference was strongly marked, and no satisfactory finishing point could be obtained. This was much more marked with the hydrochloric acid.
~Effect of Foreign Salts.~--The process for getting the arsenic into solution will exclude all metals except tin, but the solution will be charged with sodium or ammonium salts in the process of neutralising, so that it is only necessary to see if these cause any interference. The alkaline hydrates, including ammonia, are plainly inadmissible, since no free iodine can exist in their presence. Monocarbonates similarly interfere, but to a much less extent; hence the necessity for rendering the assay distinctly acid before adding the bicarbonate of soda.
With 20 c.c. of arsenic solution; and with bulk, soda, and starch as before, the results obtained were:--
"Iodine" required. With 20 grams of ammonic chloride 20.0 c.c. " 20 grams of sodium chloride 20.0 " " 20 grams of sodium acetate 20.0 " " 0.050 gram of tin, as stannic chloride 19.6 " Without any addition 20.0 "
The interference of the stannic salt is probably mechanical, the precipitate carrying down some arsenious acid.
~Effect of Varying Arsenic.~--With bulk, starch, and soda as before, but with varying arsenic, the results were:--
Arsenic added 1.0 c.c. 10.0 c.c. 20.0 c.c. 50.0 c.c. 100.0 c.c. "Iodine" required 1.1 " 9.9 " 20.0 " 50.0 " 100.0 "
~Determination of Arsenic in Metallic Copper.~--Put 1 gram of the copper filings, freed from particles of the file with a magnet, into a 16-oz.-flask; and distil with the ferric chloride mixture, as above described. Neutralise the distillate; acidify; add bicarbonate of soda and starch; dilute; and titrate with the standard solution of iodine.[107] Make a blank determination with 1 gram of electrotype copper, proceeding exactly as with the assay; and deduct the amount of arsenic found in this experiment from that previously obtained.
Working in this way on a copper containing 0.38 per cent. of arsenic and 0.80 per cent. of antimony, 0.38 per cent. of arsenic was found.
~Determination of White Arsenic in Crude Arsenic.~--Weigh out 1 gram of the dried and powdered substance (or 0.5 gram if rich), and digest with 10 c.c. of a 10 per cent. solution of soda; dilute to about 50 c.c., and filter. Render faintly acid with hydrochloric acid, and filter (if necessary); add 2 or 3 grams of bicarbonate of soda in solution, then 5 c.c. of starch, and titrate the cold solution with the standard solution of iodine.
The following is an example:--
1 gram of crude arsenic required 53.7 c.c. "Iodine;" 100 c.c. "Iodine" = 0.6000 gram white arsenic; 100 : 53.7 :: 0.6 : 0.3222, or 32.2 per cent.
With the test-tube method of dry assaying, this same sample gave results varying from 33 to 35 per cent. of white arsenic, which (judging from its appearance) was impure.
URANIC ACETATE PROCESS.
This may be looked upon as an alternative to the gravimetric method. It is applicable in all cases where the arsenic exists in solution as arsenic acid or as arsenate of soda. The process may be considered in two parts: (1) the preparation of the solution, and (2) the titration.
~Preparation of the Solution.~--If the arsenic has been separated as sulphide, it is sufficient to attack it with 10 or 15 c.c. of nitric acid, and to heat gently till dissolved, avoiding too high a temperature at first. Afterwards continue the heat till the separated sulphur runs into globules, and the bulk of the acid has been reduced to 3 or 4 c.c. Dilute with 20 or 30 c.c. of water; put in a piece of litmus paper; and add dilute ammonia until just alkaline. Then add 5 c.c. of the sodium acetate and acetic acid solution (which should make the solution distinctly acid); dilute to 150 c.c., and heat to boiling. The solution is ready for titrating.
When the arsenic exists in a nitric acid solution mixed with much copper, it is separated in the way described under Examination of Commercial Copper (Arsenic and Phosphorus), pages 208, 209.
If the arsenic has been separated as ammonium-magnesium arsenate, and phosphates are known to be absent; dissolve the precipitate (after filtering, but without washing) in dilute hydrochloric acid. Add dilute ammonia till a slight precipitate is formed, and then 5 c.c. of the sodium acetate and acetic acid solution; dilute to 150 c.c., and heat to boiling. Titrate.
If phosphates are present (which will always be the case if they were present in the original substance, and no separation with sulphuretted hydrogen has been made), the phosphorus will count in the subsequent titration as arsenic (one part of phosphorus counting as 2.4 parts of arsenic). It will be necessary to dissolve the mixed arsenate and phosphate of magnesia in hydrochloric acid. Add about four or five times as much iron (as ferric chloride) as the combined phosphorus and arsenic present will unite with, and separate by the "basic acetate" process as described under PHOSPHORUS in the Examination of Commercial Copper, page 209. Obviously, when phosphates are present, it is easier to separate the arsenic as sulphide than to precipitate it with the "magnesia mixture."
~The Titration.~--The standard solution of uranium acetate is made by dissolving 34.1 grams of the salt (with the help of 25 c.c. of acetic acid) in water; and diluting to 1 litre. The water and acid are added a little at a time, and warmed till solution is effected; then cooled, and diluted to the required volume: 100 c.c. will equal 0.50 gram of arsenic.
The sodic acetate and acetic acid solution is made by dissolving 100 grams of sodic acetate in 500 c.c. of acetic acid, and diluting with water to 1 litre. Five c.c. are used for each assay.
The solution of potassic ferrocyanide used as indicator is made by dissolving 10 grams of the salt in 100 c.c. of water.
To standardise the solution of uranium acetate, weigh up a quantity of white arsenic (As_{2}O_{3}) which shall be about equivalent to the arsenic contained in the assay (0.1 or 0.2 gram); transfer to a flask, and dissolve in 10 c.c. of nitric acid with the aid of heat. Evaporate to a small bulk (taking care to avoid the presence of hydrochloric acid); dilute with water; add a small piece of litmus paper; render faintly alkaline with ammonia; then add 5 c.c. of the sodic acetate mixture; dilute to 150 c.c.; and heat to boiling.
Fill an ordinary burette with the uranium acetate solution, and run into the assay a quantity known to be insufficient. Again heat for a minute or two. Arrange a series of drops of the solution of ferrocyanide of potassium on a porcelain slab, and, with the help of a glass rod, bring a drop of the assay solution in contact with one of these. If no colour is produced, run in the uranium acetate, 1 c.c. at a time, testing after each addition, till a brown colour is developed. It is best to overdo the assay, and to count back. It is not necessary to filter off a portion of the assay before testing with the "ferrocyanide," since the precipitate (uranic arsenate) has no effect.
The following experiments show the effect of variation in the conditions of titration. Make a solution of arsenic acid by dissolving 4.95 grams of arsenious acid (As_{2}O_{3}) in a covered beaker with 35 c.c. of nitric acid; evaporate down to 7 or 8 c.c.; and dilute with water to 1 litre: 100 c.c. will contain 0.375 gram of arsenic. Use 20 c.c. for each experiment.
~Effect of Varying Temperature.~--It is generally recommended to titrate the boiling solution, since it is possible that the precipitation is only complete on boiling. Low results are obtained in a cold solution, the apparent excess of uranium acetate striking a colour at once; on boiling, however, it ceases to do so; consequently, the solution should always be boiled directly before testing.
In four experiments made in the way described, but with 20 c.c. of a solution of arsenic acid stronger than that given (100 c.c. = 0.5 gram As), the results at varying temperatures were:--
Temperature 15° C. 30° C. 70° C. 100° C. "Uranium" required 18.0 c.c. 18.5 c.c. 18.5 c.c. 18.7 c.c.
~Effect of Varying Bulk.~--These experiments were like those last mentioned, but were titrated boiling, and the volume was varied:--
Bulk 50.0 c.c. 100.0 c.c. 200.0 c.c. 300.0 c.c. "Uranium" required 14.0 " 14.0 " 14.5 " 15.0 "
Considerable variations in bulk are to be avoided.
~Effect of Varying Sodium Acetate.~--These experiments were carried out like those last noticed, but the bulk was 150 c.c., and varying amounts of sodic acetate were added in excess of the quantity used in the experiments previously described:--
Sodic acetate added 0 gram 1 gram 10 grams 20 grams "Uranium" required 14.5 c.c. 14.5 c.c. 16.0 c.c. 18.0 c.c.
It is evidently important that the quantity of this salt present in each titration be measured out, so as to avoid variation.
~Effect of Varying the Sodium Acetate and Acetic Acid Solution.~--Acetic Acid also affects the results, but in the opposite direction, by preventing the precipitation of uranium arsenate. With varying volumes of the solution now under notice, the results were:--
Solution added 0.0 c.c. 5.0 c.c. 10.0 c.c. 15.0 c.c. "Uranium" required 14.5 " 14.5 " 14.5 " 14.0 " Solution added 20.0 " 30.0 " 40.0 " 50.0 " "Uranium" required 13.2 " 10.0 " 6.0 " 2.0 "
These show that the quantity ordered (5 c.c.) must be adhered to.
~Effect of Foreign Salts.~--In these experiments, 10 grams of the salt (the effect of which it was desired to determine) were added to a solution in other respects resembling those previously used:--
Salt added {Ammonic Ammonic Ammonic Magnesium {sulphate nitrate chloride sulphate "Uranium" required 15.5 c.c. 15.5 c.c. 15.3 c.c. 15.3 c.c.
Without any addition, 15.0 c.c. were required; and in another experiment, in which 30 grams of ammonic salts were present, 15.6 c.c. of uranium solution were required. Such variations in the amount of ammonic salts as occur in ordinary working are unimportant.
Phosphates, of course, interfere. In fact, the uranium acetate solution can be standardised by titrating with a known weight of phosphate, and calculating its equivalent of arsenic. Thus, in an experiment with 0.6 gram of hydric sodic phosphate (Na_{2}HPO_{4}.12H_{2}O), equivalent to 0.05195 gram of phosphorus, or 0.1256 gram of arsenic, 23.25 c.c. of a solution of uranium acetate were required. The same solution standardised with white arsenic gave a standard of which 100 c.c. = 0.5333 gram arsenic. On this standard the 0.6 gram of sodic phosphate should have required 23.5 c.c.
Experiments in which 0.1 gram of bismuth and 0.1 gram of antimony were present with 0.1 gram of arsenic, showed no interference on the titration. Ferric or aluminic salts would remove their equivalent of arsenic, and, consequently, must be removed before titrating.
~Effect of Varying Arsenic.~--Varying amounts of metallic arsenic were weighed up and dissolved in nitric acid, &c., and titrated:--
Arsenic taken 0.010 gram 0.050 gram 0.100 gram 0.200 gram Arsenic found 0.010 " 0.050 " 0.100 " 0.197 "
These experiments show that the method yields good results within these limits.
~Determination of Arsenic in Mispickel.~--Weigh up 1 gram of the dried and powdered ore, and evaporate to near dryness with 20 c.c. of dilute nitric acid. Make up to 100 c.c. with water, and pass sulphuretted hydrogen to reduce the ferric iron to the ferrous state, then add 20 c.c. of dilute ammonia, and again pass sulphuretted hydrogen. Warm, filter, and evaporate the filtrate to drive off the excess of ammonia; then add 10 c.c. of nitric acid, and boil down till the sulphide of arsenic at first precipitated is dissolved; neutralise; add 5 c.c. of sodium acetate and acetic acid solution; transfer to a pint flask, boil, and titrate.
For example, an impure sample of ore required, in duplicate assay of half a gram each, when treated in the above-mentioned way, 39.6 and 39.5 c.c. of the uranium acetate solution (100 c.c. = 0.537 gram of arsenic), equivalent to 0.2114 gram of arsenic, or 42.3 per cent.
An alternative method is as follows. Powder the ore very finely and weigh up .5 gram. Place in a 2-3/4 inch berlin dish and add strong nitric acid, one drop at a time until the action ceases; with care there need be no very violent reaction. Dry over a water bath. Cover with 2 grams of nitre and over this spread 5 grams of a mixture of equal parts of nitre and carbonate of soda. Fuse in a muffle or over a large gentle blow-pipe flame for 4 or 5 minutes. This will spoil the dish. Allow to cool and boil out in a larger dish with 100 c.c. of water. Filter and wash into an 8 oz. flask. Acidify the liquor with nitric and boil down to about 100 c.c. The acid should not be in too large excess, but an excess is needed to destroy nitrites. Neutralise with soda or ammonia. Add 5 c.c. of the mixture of sodium acetate and acetic acid. Titrate with uranium acetate.
~Determination of Arsenic (As) in Crude Arsenic.~--The method given under the iodine titration simply determines that portion of the arsenic which is present in the substance as arsenious oxide or white arsenic. The following method will give the total arsenic in the sample. It would be incorrect to report this as so much per cent. of arsenious oxide, although it may be reported as so much per cent. of arsenic equivalent to so much per cent. of white arsenic, thus:--
Arsenic 30.0 per cent. Equivalent to white arsenic 39.6 "
The equivalent of white arsenic is calculated by multiplying the percentage of arsenic by 1.32. The method of determining the percentage of arsenic is as follows:---Boil 1 gram of the sample with 10 c.c. of nitric acid. When the bulk of the solution has been reduced to one-half, and red fumes are no longer evolved, dilute with a little water, and filter into a flask. Neutralise the filtrate, add 5 c.c. of sodic acetate solution, boil and filter. The precipitate (ferric arsenate) is transferred to a small beaker, treated with 5 c.c. of dilute ammonia, and sulphuretted hydrogen passed through it. The iron sulphide is filtered off, and the filtrate evaporated with an excess of nitric acid. When the solution is clear, it is neutralised, and 1 or 2 c.c. of sodic acetate solution having been added, is then mixed with the first filtrate. The solution is boiled and titrated.
A sample treated in this way required 49.2 c.c. of the uranium acetate solution (100 c.c. = 0.537 gram of arsenic), equivalent to 26.4 per cent.
~Determination of Arsenic in Brimstone.~--Take 10 grams of the substance, and powder in a mortar; rub up with 10 c.c. of dilute ammonia and a little water; rinse into a pint flask; pass a current of sulphuretted hydrogen; and warm on a hot plate for a few minutes. Filter, acidulate the filtrate with sulphuric acid; filter off the precipitate; attack it with 10 c.c. of nitric acid; and proceed as in the other determinations.
PRACTICAL EXERCISES.
1. Mispickel contains 45.0 per cent. of arsenic, to how much white arsenic will this be equivalent?
2. How would you make a standard solution of iodine so that 100 c.c. shall be equivalent to 1 gram of white arsenic?
3. What weight of arsenic is contained in 1 gram of pyrarsenate of magnesia, and what weight of ammonic-magnesic arsenate would it be equivalent to?
4. The residue, after heating 10 grams of crude arsenic, weighed 0.62 gram. What information does this give as to the composition of the substance? If another 10 grams of the substance, heated on a water-bath, lost 0.43 gram, what conclusions would you draw, and how would you report your results?
5. If a sample of copper contained 0.5 per cent. of arsenic, and 1 gram of it were taken for an assay, how much standard uranium acetate solution would be required in the titration?
PHOSPHORUS AND PHOSPHATES.
Phosphorus rarely occurs among minerals except in its highest oxidized state, phosphoric oxide (P_{2}O_{5}), in which it occurs abundantly as "rock phosphate," a variety of apatite which is mainly phosphate of lime. Phosphates of most of the metallic oxides are found. Phosphoric oxide in small quantities is widely diffused, and is a constituent of most rocks. Its presence in varying amounts in iron ores is a matter of importance, since it affects the quality of the iron obtainable from them.
Phosphorus occurs in alloys in the unoxidized state. It is directly combined with the metal, forming a phosphide. In this manner it occurs in meteoric iron. The alloy phosphor-bronze is made up of copper, tin, zinc, and phosphorus.
Phosphates are mined in large quantities for the use of manure manufacturers, and for making phosphorus.
Phosphorus and arsenic closely resemble each other in their chemical properties, more especially those which the assayer makes use of for their determination. Phosphorus forms several series of salts; but the phosphates are the only ones which need be considered. Pyrophosphate of magnesia, which is the form in which phosphoric oxide is generally weighed, differs from the ordinary phosphate in the proportion of base to acid. Metaphosphates differ in the same way. If these are present, it must be remembered they act differently with some reagents from the ordinary phosphates, which are called orthophosphates. They are, however, all convertible into orthophosphates by some means which will remove their base, such as fusion with alkaline carbonates, boiling with strong acids, &c.[108]
Phosphides are converted into phosphates by the action of nitric acid or other oxidizing agents. Dilute acids, when they act on the substance, evolve phosphuretted hydrogen (PH_{3}). The student should be on his guard against losing phosphorus in this manner.
There is no dry assay for phosphorus. All assays for it are made either gravimetrically or volumetrically.
The separation of phosphoric oxide is made as follows:--The ore or metal is dissolved in acid and evaporated, to render the silica insoluble. It is taken up with hydrochloric acid, diluted with water, and treated with sulphuretted hydrogen. The filtrate is boiled, to get rid of the excess of gas, and treated with nitric acid, to peroxidize the iron present. If the iron is not present in more than sufficient quantity to form ferric phosphate with all the phosphorus present, some ferric chloride is added. The iron is then separated as basic acetate. The precipitate will contain the phosphorus, together with any arsenic acid not reduced by the sulphuretted hydrogen. The precipitate should have a decided brown colour. The precipitate is washed, transferred to a flask, and treated first with ammonia, and then with a current of sulphuretted hydrogen. The filtrate from this (acidulated with hydrochloric acid, and, if necessary, filtered) contains the phosphorus as phosphoric acid. This method is not applicable in the presence of alumina, chromium, titanium, or tin, if the solution is effected with nitric acid. The precipitate obtained by the action of nitric acid on tin retains any phosphoric or arsenic oxide that may be present.
A method of separation more generally applicable and more convenient to work is based on the precipitation of a yellow phospho-molybdate of ammonia,[109] by the action of an excess of ammonic molybdate upon a solution of a phosphate in nitric acid. Dissolve the substance by treatment with acid, and evaporate to dryness. Take up with 10 c.c. of nitric acid, and add 20 grams of ammonic nitrate, together with a little water. Next put in the solution of ammonium molybdate solution in the proportion of about 50 c.c. for each 0.1 gram of phosphoric oxide judged to be present. Warm to about 80° C., and allow to stand for an hour. Filter, and wash with a 10 per cent. solution of ammonic nitrate. It is not necessary that the whole of the precipitate be placed on the filter; but the beaker must be completely cleaned. Dissolve the precipitate off the filter with dilute ammonia, and run the solution into the original beaker. Run in from a burette, slowly and with stirring, "magnesia mixture," using about 15 c.c. for each 0.1 gram of phosphoric oxide. Allow to stand for one hour. The white crystalline precipitate contains the phosphorus as ammonium-magnesium phosphate.
Phosphate of lead is decomposed by sulphuric acid; the lead is converted into the insoluble lead sulphate, and the phosphoric acid is dissolved. Phosphate of copper and phosphate of iron may be treated with sulphuretted hydrogen; the former in an acid, and the latter in an alkaline, solution. Phosphate of alumina is generally weighed without separation of the alumina, since this requires a fusion. In all cases the aim is to get the phosphoric oxide either free, or combined with some metal whose phosphate is soluble in ammonia.
Joulie's method of separation is as follows:--One to ten grams of the sample are treated with hydrochloric acid, and evaporated to dryness with the addition (if any pyrites is present) of a little nitric acid. The residue is taken up with hydrochloric acid, cooled, transferred to a graduated flask, and diluted to the mark. It is then shaken up, filtered through a dry filter, and a measured portion (containing about 0.05 gram of phosphoric acid) transferred to a small beaker. Ten c.c. of a citric-acid solution of magnesia[110] is added, and then an excess of ammonia. If an immediate precipitate is formed, a fresh portion must be measured out and treated with 20 c.c. of the citrate of magnesia solution and with ammonia as before. The beaker is put aside for from two to twelve hours. The precipitate is then filtered off and washed with weak ammonia; it contains the phosphorus as ammonium-magnesium phosphate.
GRAVIMETRIC DETERMINATION.
If the phosphate is not already in the form of ammonic-magnesic phosphate, it is converted into this by the addition to its solution of an excess of ammonia and "magnesia mixture." In order to get the precipitate pure, the "magnesia mixture" is run in gradually (by drops) from a burette, with constant stirring. A white crystalline precipitate at once falls, if much phosphorus is present; but, if there is only a small quantity, it may be an hour or two before it shows itself. The solution is best allowed to rest for twelve or fifteen hours (overnight) before filtering. The presence of tartaric acid should be avoided; and the appearance of the precipitate should be crystalline. The solution is decanted through a filter, and the precipitate washed with dilute ammonia, using as little as may be necessary. The precipitate is dried, transferred to a weighed Berlin or platinum crucible; the filter-paper is carefully burnt, and its ash added to the precipitate, which is then ignited, at first gently over a Bunsen burner, and then more strongly over the blowpipe or in the muffle. The residue is a white mass of magnesium pyrophosphate containing 27.92 per cent. of phosphorus, or 63.96 per cent. of phosphoric oxide.
VOLUMETRIC METHOD.
Instead of separating and weighing this compound, the phosphoric oxide in it can be determined by titration. In many cases the ore may be dissolved and immediately titrated without previous separation. It is better, however, to carry the separation so far as to get phosphoric acid, an alkaline phosphate, or the magnesia precipitate. It may then be prepared for titration in the following way:--The precipitate in the last case (without much washing) is dissolved in a little hydrochloric acid, and the solution in any case rendered fairly acid. Dilute ammonia is added till it is just alkaline, and then 5 c.c. of the sodic acetate and acetic acid mixture (as described under the Arsenic Assay). This should yield a clear distinctly-acid solution. It is diluted to 100 or 150 c.c., heated to boiling, and titrated with the uranium acetate solution, using that of potassic ferrocyanide as indicator.
The standard solution required is made by dissolving 35 grams of uranium acetate in water with the aid of 25 c.c. of acetic acid, and diluting to 1 litre.
An equivalent solution of phosphoric oxide is made by dissolving 25.21 grams of crystallised hydric disodic phosphate (HNa_{2}PO_{4}.12H_{2}O) in water, and making up to 1 litre. 100 c.c. will contain 0.5 gram of phosphoric oxide (P_{2}O_{5}), or 0.2183 gram of phosphorus. In making this solution, transparent crystals only must be used. The uranium acetate solution is only approximately equivalent to this, so that its exact standard must be determined.
Sodic Acetate and Acetic Acid Solution.--It is the same as that described under Arsenic.[111] Use 5 c.c. for each assay.
The following experiments show the effect of variation in the conditions of the titration:--
~Effect of Varying Temperature.~--The solution should be titrated while boiling. This is especially necessary for the last few c.c. in order to get a decided and fixed finishing point.
Temperature 15° C. 30° C. 70° C. 100° C. "Uranium" required 18.0 c.c. 19.2 c.c. 19.0 c.c. 18.9 c.c.
~Effect of Varying Bulk.~--
Bulk 50.0 c.c. 100.0 c.c. 200.0 c.c. 300.0 c.c. "Uranium" required 18.8 " 18.9 " 19.0 " 19.3 "
Variation in bulk affects the results; therefore, a constant bulk should be adhered to.
~Effect of Varying Sodium Acetate and Acetic Acid Solution.~--
Sodium acetate and acetic acid solution 0.0 c.c. 1.0 c.c. 5.0 c.c. 10.0 c.c. 20.0 c.c. "Uranium" required 18.9 " 18.9 " 19.0 " 18.8 " 17.5 "
As in the titration with arsenates, an excess is dangerous to the assay; a definite quantity (5 c.c.) should, therefore, be used.
~Effect of Foreign Salts.~--Besides the sodium acetate, &c., added, the only salts likely to be present are those of ammonia and magnesia. In three experiments, in one of which no foreign salts were introduced, while in the other two 5 grams of ammonic chloride and of magnesium sulphate respectively were added, there were required:--
With ammonic chloride 18.8 c.c. "Uranium" solution With magnesium sulphate 19.0 " " Without foreign salts 18.9 " "
~Effect of Varying Phosphate.~--
"Phosphate" solution added 10.0 c.c. 20.0 c.c. 50.0 c.c. 100.0 c.c. "Uranium" required 9.8 " 18.9 " 47.6 " 94.5 "
The quantity of phosphoric oxide in the assay solution for the conditions of titration should not be much less than 0.05 gram. For smaller quantities the uranium solution should be diluted to half its strength, and the assay solution concentrated by reducing its bulk to 50 c.c. and using 2.5 c.c. of the sodium acetate and acetic acid solution.
~Determination of Phosphoric Oxide in Apatite.~--Weigh up 0.5 gram of the dried and powdered sample, and dissolve it in 5 c.c. of hydrochloric acid. Evaporate to a paste, add 5 c.c. of the sodic acetate and acetic acid solution, dilute to 100 c.c. with water, boil, and titrate with uranium acetate solution.
In an example, 0.5 gram of apatite required 37.4 c.c. of uranium acetate solution (standard equal to 0.5291 gram of phosphoric oxide). The sample therefore contained 0.1979 gram of P_{2}O_{5}, equal to 39.58 per cent.
~Determination of Phosphoric Oxide in an Iron Ore.~--Take 10 grams, boil with 50 c.c. of hydrochloric acid, and evaporate to a paste; take up with 10 c.c. of dilute hydrochloric acid, and dilute with water to 400 c.c. Pass sulphuretted hydrogen for nearly a quarter of an hour; warm, and filter. Boil off the excess of gas; cool, add ammonia till nearly neutral, and then a few drops of ferric chloride solution, and 4 or 5 grams of sodium acetate, with a drop or two of acetic acid. Boil and filter. Dissolve the precipitate in hot dilute hydrochloric acid, and add citro-magnesia mixture and ammonia; allow to stand overnight; filter, ignite, and weigh.
In an example, 10 grams of ore gave 28.5 milligrams of magnesic pyrophosphate, which is equivalent to 0.18 per cent. of phosphoric oxide.
~Determination of Phosphorus in Iron.~--Take from 2 to 10 grams (according to the amount of phosphorus present), and dissolve in aqua regia, keeping the nitric acid in excess; evaporate to dryness and take up with hydrochloric acid, boil, dilute, and filter. Add 10 c.c. of nitric acid, nearly neutralise with ammonia, render acid with 3 or 4 c.c. of nitric acid, and add 10 or 20 c.c. of ammonic molybdate solution. Heat for some time, allow to settle, filter, and wash the precipitate with a solution of ammonic nitrate. Dissolve the precipitate in dilute ammonia, nearly neutralise with dilute hydrochloric acid, and add first "magnesia mixture," and then ammonia; allow to stand overnight; filter, wash with dilute ammonia, dry, ignite, and weigh as magnesic pyrophosphate. Calculate to phosphorus.
PRACTICAL EXERCISES.
1. Ten grams of an iron yielded 12 milligrams of pyrophosphate of magnesia. What percentage of phosphorus did the metal contain?
2. Ten grams of an iron ore gave 12 milligrams of pyrophosphate. What percentage of phosphoric oxide did it contain?
3. What weight of apatite 3Ca_{3}(PO_{4}){2}.CaClF would require 50 c.c. of standard uranium solution (100 c.c. equal to 0.5 gram of P{2}O_{5})?
4. You have reason to believe that a precipitate which has been weighed as magnetic pyrophosphate contains some arsenate. How would you determine the amount of phosphate really present?
5. Twenty c.c. of a solution of sodic phosphate containing 0.100 gram of P_{2}O_{5} was found to require a solution containing 0.700 gram of hydrated uranium acetate in a titration. The precipitate contains 80.09 per cent. uranium oxide and 19.91 per cent. of phosphoric oxide. What percentage of uranium oxide was contained in the uranic acetate?
NITROGEN AND NITRATES.
Nitrogen occurs in nature in the free state, and forms about four-fifths of the atmosphere. In combination, as nitrate, it is found in nitre (KNO_{3}), and Chili saltpetre (NaNO_{3}), minerals which have a commercial importance. The latter occurs in beds, and is extensively worked for use as a manure and in the preparation of nitric acid.
Nitrogen is mainly characterised by negative properties, although many of its compounds are very energetic bodies. It is a gas, present everywhere, but so inactive that the assayer can always afford to ignore its presence, and, except in testing furnace gases, &c., he is never called on to determine its quantity.
The nitrates are an important class of salts, and may be looked on as compounds of the bases with nitric pentoxide (N_{2}O_{5}). They are, with the exception of a few basic compounds, soluble in water, and are remarkable for the ease with which they give up their oxygen. The alkaline nitrates fuse readily, and lose oxygen with effervescence forming nitrites; while at a higher temperature they yield more oxygen and lose their nitrogen, either as a lower oxide or as nitrogen. The nitrates of the metals, on heating, leave the oxide of the metal. It is as yielders of oxygen that nitrates are so largely used in the manufacture of explosives. Gunpowder contains from 65 to 75 per cent. of potassium nitrate (nitre).
Nitrates are best detected and determined by their yielding nitric oxide when treated with sulphuric acid and a suitable reducing agent, such as ferrous sulphate, mercury, or copper. Nitric oxide is a colourless gas very slightly soluble in water. It combines at once with oxygen, on mixing with the air, to form brown "nitrous fumes," and dissolves in a solution of ferrous sulphate, producing a characteristic blackish-brown colour. It is this colour which affords the best and most easily-applied test for nitrates. The substance suspected to contain nitrates is dissolved in about 1 c.c. of water, and treated with an equal volume of strong sulphuric acid. After cooling, a solution of ferrous sulphate is poured on its surface, so as to form a layer resting on it. On standing, a brown or black ring is developed where the liquids join, if any nitrate or nitrite is present. Nitrites are distinguished from nitrates by effervescing and yielding brown fumes when treated with a little dilute sulphuric acid.
The separation of nitrates is in many cases difficult. Generally, on treating the substance with water, the nitrate will be in the solution, and is filtered off from any insoluble matter. In the exceptional cases it is got into solution by treating with a boiling solution of sodium carbonate; the nitrate will contain it as an alkaline nitrate.
Since, however, in their determination, nitrates are never separated and weighed as such, the difficulty of separating them has little importance. Usually, the determination can be made on the original aqueous solution, and it is never necessary to do more than remove any special substance which has a bad effect; and this is easily done by the usual reagents.
GRAVIMETRIC DETERMINATION.
It follows from what has been said that there is no direct gravimetric determination. The percentage of nitrogen pentoxide (N_{2}O_{5}) in a comparatively pure nitrate is sometimes determined indirectly in the following way:--Place in a platinum-crucible 4 or 5 grams of powdered and cleaned quartz. Ignite, cool in a desiccator, and weigh with the cover. Mix 1 gram of the dried and powdered salt with the quartz in the crucible by stirring with a stout platinum-wire. Cover the crucible, and heat in a Bunsen-burner flame at scarcely visible redness for half-an-hour. Cool and weigh. The loss in weight gives the amount of nitrogen pentoxide. Sulphates and chlorides in moderate quantity do not interfere. The following is an example of the process:--
Crucible and sand 26.6485 grams Nitre taken 1.0000 " ------- 27.6485 " Weight after ignition 27.1160 " ------- Loss on ignition 0.5325 "
This is equal to 53.25 per cent. of nitrogen pentoxide.
VOLUMETRIC DETERMINATION.
This is based on the oxidising action of nitric acid, or of nitrates in acid solutions on ferrous salts. The pentoxide (N_{2}O_{5}) of the nitrate is reduced to nitric oxide (NO), so that 336 parts of iron peroxidised represent 108 parts of nitric pentoxide as oxidising agent.[112] The quantity of iron peroxidised is determined by taking a known quantity of ferrous salt, oxidizing with a weighed sample of nitrate, and then determining the residual ferrous iron by titration with bichromate or permanganate of potassium solution. The difference between the ferrous iron taken and that found, gives the amount oxidized by the nitrate. The speed with which nitric oxide takes up oxygen from the air, and thus becomes capable of oxidising more iron, renders some precautions necessary; ferrous chloride should, therefore, be used, since it is easier to expel nitric oxide (by boiling) from solutions of a chloride than it is from those of a sulphate. The process is as follows:--Dissolve 2 grams of thin soft iron wire in 50 c.c. of hydrochloric acid in a flask provided with an arrangement for maintaining an atmosphere of carbon dioxide. When the iron has dissolved, allow the solution to cool, and add 0.5 gram of the nitrate. Heat gently for a few minutes, and then boil until the nitric oxide is expelled. An atmosphere of carbon dioxide must be kept up. Dilute with water, and titrate the residual iron with standard solution of bichromate of potassium. The standard "bichromate" is made by dissolving 17.5 grams of the salt (K_{2}Cr_{2}O_{7}) in water, and diluting to 1 litre: 100 c.c. equal 2 grams of iron. Deduct the weight of iron found from the 2 grams originally taken, and multiply by 0.3214. This gives the weight of the pentoxide in the sample. In an example, 0.5 gram of nitre was taken, and 59.4 c.c. of the "bichromate" solution were required. The 59.4 c.c. thus used are equivalent to 1.198 gram of iron. This leaves 0.822 gram as the quantity oxidised by the nitre, which, multiplied by 0.3214, gives 0.2642 gram for the nitrogen pentoxide, or 52.8 per cent.
GASOMETRIC METHOD.
This is based upon the measurement of the nitric oxide evolved on shaking up a weighed quantity of the nitrate with sulphuric acid over mercury in a nitrometer. Each c.c. of nitric oxide obtained, when reduced to normal temperature and pressure, is equivalent to:--
0.627 milligram of nitrogen. 1.343 " of nitric oxide. 2.418 " of nitric pentoxide. 2.820 " of nitric acid. 3.805 " of sodium nitrate. 4.523 " of potassium nitrate.
In working on substances not rich in nitrates, an ordinary nitrometer (fig. 69) is used; but in the assay of sodium nitrate, nitroglycerine, &c., an instrument provided with a bulb having a capacity of 100 c.c. is employed.
The plan of working is as follows:--The "measuring tube" is filled with mercury until it reaches up into the tap, and the levelling-tube is placed so that it contains an inch or two of mercury. If the nitrate is in solution, 2 or 3 c.c. of the liquid (dilute liquids are brought to this bulk by evaporation) are measured into the cup. The levelling-tube is lowered a little, and the tap cautiously opened until all but the last drop of the liquid has run in. The cup is then rinsed with 2 or 3 c.c. of sulphuric acid, which is run in in the same way, and the operation is repeated with another lot of acid. The measuring-tube is now taken from the clamp, and shaken for two or three minutes, until no more gas is given off. It is replaced, and the mercury-level in the two tubes adjusted. Then it is allowed to stand until the froth has subsided, and the gas has cooled to the temperature of the room. The volume of the gas is then read off. In adjusting the level, account must be taken of the sulphuric acid in the measuring-tube; this is allowed for by having the mercury higher in the other tube by, say, 1 mm. for each 6.5 mm. of sulphuric acid, or it is counterpoised by an equal height of sulphuric acid in the levelling-tube, in which case the two mercury-levels are made to correspond. On opening the tap after reading off the volume, there should be no change in the level of the mercury. If it should rise or fall a little, a slight increase or decrease (say 0.1 c.c.) is made to the volume previously read off.
In working with nitrate of soda, &c., in the bulb nitrometer, it is necessary to take a quantity of the substance which will yield more than 100 and less than 150 c.c. of the gas.
FOOTNOTES:
[103] Na_{3}AsO_{3} + H_{2}O + 2I = Na_{3}AsO_{4} + 2HI. The acid is at once neutralised.
[104] Mr. Thomas Gibb is the originator of this ingenious process.
[105] By taking hold of the water present, it may prevent the dissociation of arsenious chloride.
[106] It is difficult to get ferric chloride free from arsenic; but the following treatment will remove 80 or 90 per cent. of the arsenic contained in the commercial material:--Dissolve 2 or 3 lbs. of ferric chloride with the smallest amount of water that will effect solution with the addition of 100 c.c. of hydrochloric acid; add a solution of sulphurous acid in quantity sufficient to reduce 2 or 3 per cent. of the iron to the ferrous state; allow to stand a week; and then boil, to remove the hydrochloric acid added. Nitric acid, which is prejudicial, is also removed by this treatment.
[107] When the amount of arsenic to be estimated is small (as in refined coppers), it is better to use a weaker solution of iodine. This is made by diluting 200 c.c. of the standard solution with water to 1 litre. Each c.c. will equal 0.1 per cent., if 1 gram of the metal has been taken for the assay.
[108] The constitution of these phosphates may be thus illustrated--
Magnesic meta-phosphate MgO.P_{2}O_{5}. Magnesic pyro-phosphate 2MgO.P_{2}O_{5}. Magnesic ortho-phosphate 3MgO.P_{2}O_{5}.
[109] The composition of which is--
MoO_{2} 90.74, P_{2}O_{5} 3.14, (NH_{4})_{2}O 3.57, H_{2}O 2.55 = 100.00.
[110] This is made by adding 27 grams of magnesium carbonate (a little at a time) to a solution of 270 grams of citric acid in 350 c.c. of warm water; and, when dissolved, adding 400 c.c. of dilute ammonia, and making up the bulk to 1 litre; 20 c.c. of the solution is sufficient for 0.1 gram of P_{2}O_{5}, although more will be required if much iron or alumina is present.
[111] For the details of the titration, the student is referred to the same place.
[112] N_{2}O_{5} + 6FeO = 3Fe_{2}O_{3} + 2NO.