Tungstic acid50.1%Cassiterite10.9Ferrous oxide24.6Manganous oxide5.4Niobic oxide, alumina, &c.3.5Silica1.2Copper oxide2.7Zinc oxide0.22Arsenic0.51Sulphur0.20———99.33
These oxides are commonly met with in samples of wolfram and tinstone, especially niobic. They are probably present in the form of columbite, a niobate of iron and manganese; and tantalite, a tantalate of the same metals.
On boiling with hydrochloric acid they are both liberated, and remain for the greater part (all the niobic) in the insoluble residue with the tungstic acid. On removing the latter with dilute ammonia they remain as a white insoluble precipitate, very prone to run through the filter on washing. They may be dissolved in hydrofluoric acid either at once or after fusion with bisulphate of potash, and extraction with cold water. To the solution in hydrofluoric acid gradually add a boiling solution of acid potassium fluoride (HF, KF.). Potassic fluotantalate (soluble in 200 parts of water) separates out first, and afterwards potassic fluoniobate (soluble in 12 parts of water). The separated salts (after heating with sulphuric acid and washing out the potassium sulphate formed) are ignited with ammonic carbonate, and weighed as tantalic oxide (Ta2O5) and niobic oxide (Nb2O5) respectively.
They are both white powders. The oxide of niobium dissolved in a bead of microcosmic salt gives a bluish colour in the reducing flame. The oxide of tantalum dissolves in the bead, but gives no colour.
FOOTNOTES:[76]This will give almost the whole of the tin; a further portion will be got in subsequent work, and must be added to this result.[77]Published by P. Holland, in theChemical News, vol. lix. p. 27.
[76]This will give almost the whole of the tin; a further portion will be got in subsequent work, and must be added to this result.
[76]This will give almost the whole of the tin; a further portion will be got in subsequent work, and must be added to this result.
[77]Published by P. Holland, in theChemical News, vol. lix. p. 27.
[77]Published by P. Holland, in theChemical News, vol. lix. p. 27.
Manganese occurs mainly as black oxide (MnO2) in the mineral pyrolusite; and, in a less pure form, in psilomelane and wad. The value of the ore depends rather on the percentage of available oxygen than on the proportion of metal present. The results of assays are generally reported as so much per cent. of the dioxide (MnO2). In smaller quantities it is very widely distributed. Manganese itself has a value for steel-making; or, rather, for the making of spiegeleisen and ferro-manganese, which are used in the Bessemer and Siemens processes. For this purpose the percentage of the metal (Mn) is required. Consequently the minerals of manganese may be considered in two aspects—(1) as a source of oxygen; and (2) as a source of manganese. These will require separate consideration.
The black oxide is mainly used in the preparation of chlorine, liberation of which it brings about when treated with hot hydrochloric acid, or with a mixture of common salt and sulphuric acid. The quantity of chlorine which is obtained depends upon the proportion of dioxide present;[78]and in assaying may either be measured by its equivalent of iodine liberated, or by the oxidising effect on an acid solution of ferrous sulphate. When the ore also carries substances which have a reducing effect (such as ferrous compounds), such assays will give, not the total dioxide (MnO2), but less, by the amount required to oxidise these impurities; and this is exactly what is required in valuing such an ore for commercial purposes. Manganese compounds are characterised by the readiness with which they may be converted into highly-oxidised bodies. Solution of manganese in hydrochloric acid, rendered alkaline with ammonia, yields a clear solution,[79]which rapidly takes up oxygen from the air, forming a brown precipitate of the oxide (Mn2O3). The addition ofbromine or chlorine to such a solution determines the precipitation of a still higher oxide (approximately MnO2). On treating a compound containing manganese with nitric acid and dioxide of lead (PbO2), the oxidation is carried still further, a purple-coloured solution of permanganic acid (HMnO4or H2O.Mn2O7) being formed. On fusing minerals containing (even traces of) manganese with sodium carbonate in an open crucible, a green "melt" is obtained which owes its colour to sodium manganate (Na2MnO4or Na2O.MnO3). This salt is soluble in water, forming a green solution; which, when rendered acid, rapidly changes into the permanganate with the characteristic purple colour. Permanganate of potash is a salt much used in assaying, with some properties of which the student will have already become familiar.
Compounds of manganese, on boiling with strong hydrochloric acid, yield manganous chloride[80](MnCl2).
The properties given above serve for the detection of manganese; the higher oxides are distinguished by causing the evolution of chlorine (with its peculiarly suffocating smell) when acted on with hydrochloric acid; while the green "melt," with sodium carbonate, can be relied on for the recognition of manganese itself. There is no dry assay of manganese ores.
Strong hydrochloric acid is the best solvent for ores of manganese; but where the proportion of dioxide (MnO2) is required, the solution is effected during the assay. The ore should be in a very fine state of division before treatment with acids.
The separation of manganese from other metals is thus effected: Ignite, in order to destroy any organic matter which may be present; dissolve in hydrochloric acid, and evaporate to dryness, to separate silica. Take up with hydrochloric acid, dilute, pass sulphuretted hydrogen, and filter. Boil off the excess of gas, peroxidise the iron with a drop or two of nitric acid, and separate the iron as basic acetate (as described underIron).[81]If the iron precipitate is bulky, it is dissolved in a little hydrochloric acid, reprecipitated, and the filtrate added to the original one. Neutralise with soda, and add bromine in excess; heat gradually to boiling, allow to settle, and filter. The precipitate is impure dioxide of manganese (containing alkalies and, possibly, cobalt or nickel).
Dissolve the precipitate in hydrochloric acid, and boil; add a slight excess of carbonate of soda, warm, and filter. Wash with hot water, dry, carefully ignite in an open Berlin crucible, and weigh. The substance is the brown oxide (Mn3O4), and contains 72.05 per cent. of manganese. If the percentage of dioxide is required it may be calculated by multiplying the percentage of manganese by 1.582. It must be borne in mind that the manganese should never be calculated to dioxide except when it is known to exist in the ore only in that form.
The two methods are based on the oxidising effect of manganese dioxide; and if the metal does not already exist in this form it will require a preliminary treatment to convert it. The following method due to Mr. J. Pattinson[82]effects this: A quantity of the ore containing not more than .25 grams of the metal (Mn), is dissolved in hydrochloric acid in a pint beaker, and, if necessary, 3 or 4 c.c. of nitric acid are added to peroxidise the iron, and ferric chloride is added if required, so that there may be at least as much iron as manganese. Calcium carbonate is added till the solution is slightly red; and next the redness is removed by the cautious addition of acid; 30 c.c. of zinc chloride solution (containing 15 grams of zinc per litre) are added, the liquid is brought to boil and diluted to about 300 c.c. with boiling water; 60 c.c. of a solution of bleaching powder (33 grams to the litre and filtered), rendered slightly greenish by acid, are then run in and are followed by 3 grams of calcium carbonate suspended in 15 c.c. of boiling water. During effervescence the beaker is covered, the precipitate is stirred, and 2 c.c. of methylated spirit are mixed in. The precipitate is collected on a large filter, washed with cold water, and then with hot, till free from chlorine, which is tested for with starch and potassium iodide. The acid ferrous sulphate solution (presently described) is then measured into the beaker, and the precipitate, still in the paper, added; more acid is added (if necessary), and the solution is diluted and titrated. In place of bleaching powder solution, 90 c.c. of bromine water (containing 22 grams per litre) may be used.
This method, which is the one commonly used, is based on the determination of the amount of ferrous iron oxidised by a known weight of the ore. It is known that 87 parts of the dioxide will oxidise 112 parts of ferrous iron;[83]therefore 1 gram will oxidise 1.287 gram of ferrous iron, or 1 gram of ferrous iron oxidised will be equivalent to 0.7768 gram of the dioxide. The finely-divided substance containing the dioxide is digested in a solution of a known quantity of iron in sulphuric acid. The iron, of course, must be in excess, which excess is determined when the ore is dissolved by titrating with standard permanganate or bichromate of potash solution. The assay resolves itself into one for the determination of ferrous iron, for which the standard solutions and method of working described underIronare used.
The assay is as follows:—For rich ores, 2 grams of clean soft iron wire are treated, in a pint flask, with 100 c.c. of dilute sulphuric acid and warmed till dissolved. Carefully sample the ore, and in one portion determine the "moisture at 100° C.;" grind the rest in a Wedgwood mortar with a little pure alcohol until free from grit. This reduces the substance to a finely-divided state and assists solution. Evaporate off the alcohol and dry at 100° C., mix well, and keep in a weighing-bottle. Weigh up 2 grams and add them to the solution of iron in the flask; carefully wash it all down into the acid liquid. On rotating the flask the ore will rapidly dissolve, but gentle heat may be used towards the end to complete the solution. When the residue is clean and sandy-looking, and free from black particles, the flask is cooled, and the residual ferrous iron is determined by titration with "permanganate." The iron thus found, deducted from the 2 grams taken, will give the amount of iron peroxidised by the dioxide contained in the 2 grams of ore. This divided by 2 and multiplied by 77.68 will give the percentage of dioxide in the sample, or multiplied by 49.41 will give that of metallic manganese.
When the quantity of manganese or of the dioxide to be determined is small, it is not necessary to use 2 grams of iron; 1 gram, or even less, may be taken. The iron may be used in the form of a standard solution of ferrous sulphate and portions measured off, thus saving the labour of weighing.
Determination of Dioxide in a Manganese Ore.—Weigh up 1 or 2 grams of the finely-powdered ore[84]and an equal weight of pure iron wire, dissolve the wire in 50 or 100 c.c. of dilute sulphuric acid, and, when solution is complete, add the ore and warm till it too is dissolved. Cool and titrate the remaining ferrous iron with the permanganate or bichromate of potassium solution.
For example, 0.7560 gram of pyrolusite and 1.000 gram of iron were taken and treated as above; 13.9 c.c. of "permanganate" (standard 100 c.c. = 0.4920 gram iron) were required; this indicates that 0.0684 gram of iron was left unoxidised by the ore. The iron oxidised, then, was 0.9316 gram (1.000 - 0.0684); multiplying this by 0.7768, we find that 0.7237 gram is the quantity of manganese dioxide which was present. This is equivalent to 95.77 per cent.;
0.7560 : 0.7237 :: 100 : 95.77.
It has been already stated that when dioxide of manganese is boiled with strong hydrochloric acid chlorine is given off, and that the amount of chlorine so liberated is a measure of the dioxide present. If the chlorine is passed into a solution of potassium iodide, an equivalent of iodine will be set free.[85]This is apparently a very indirect way of determining how much of the dioxide is present; but the reactions are very sharp, and the final determination of the iodine is an easy one.
Fig. 60.
The finely-powdered sample of dioxide is placed in a small flask provided with an exit tube leading into a solution of potassic iodide (fig. 60). On adding hydrochloric acid and boiling, the chlorine evolved is driven into the iodide solution and there absorbed; the boiling is continued till the steam and hydrochloric acid fumes have driven the last portions of the chlorine out of the flask and into the solution. In this experiment there is a strong tendency for the iodide solutionto rush back into the flask. This tendency is overcome by avoiding draughts and regulating the heat; or by placing a lump of magnesite in the flask, which acts by evolving carbonic acid and so producing a steady outward pressure. When the distillation is finished the tube containing the iodine is detached and washed out into a beaker. If the solution is strongly acid it should be almost neutralised by the cautious addition of dilute ammonia. If crystals of iodine have separated, potassium iodide must be added in quantity sufficient to dissolve them. The condenser must be kept cool whilst the chlorine is passing into it.
The solution, transferred to a beaker, is titrated with a standard solution of sodic hyposulphite (100 c.c. = 1.27 gram iodine or 0.435 gram of dioxide of manganese). In titrating, the solution should be cold, or not warmer than 30° C. The bulk may vary from 100 to 200 c.c.; but it is best always to work with the same volume. The "hypo" is run in with constant agitation until the brown colour has been reduced to a light yellow; 5 c.c. of starch solution are then added and the titration cautiously continued until the end is reached; the finish is indicated by a change from blue to colourless.
The assay solution may be acidified with acetic, sulphuric, or hydrochloric acid before titrating with "hypo;" but it must be only faintly so. An excess of acid may be nearly neutralised with ammonia without interference, but excess of alkali is fatal. Bicarbonate of soda must not be used in excess; it is best to avoid it altogether. The assay solution should be titrated at once, as it weakens on standing; and the "hypo" solution should be standardised every two or three days, as its strength is not constant.
The standard solution of hyposulphite of sodais made by dissolving 25 grams of the salt (Na2S2O3.5H2O) in water and diluting to 1 litre. 100 c.c. are equivalent to 1.27 gram of iodine.
This solution is standardised by weighing, in a small beaker, about half a gram of iodine, to which is added a crystal or two of potassium iodide and a few drops of water. When dissolved, the solution is diluted to 100 c.c., and titrated in the manner described. The starch solution is made in the manner described under the iodide copper assay. 5 c.c. are used for each titration.
In determining the effects of variations in the condition of the assay a solution of iodine was used, which was equivalent in strength to the "hypo" solution. It was made by dissolving 12.7 grams of iodine with 25 grams of potassium iodide in a little water and diluting to 1 litre. 100 c.c. of this solution were found (at the time of the experiments) to be equivalent to 102.0 c.c. of the "hypo."
Effect of Varying Temperature.—The bulk of the solution was 100 c.c.; 20 c.c. of iodine were taken, and 5 c.c. of starch solution were added towards the end as indicator. These conditions are also those of the other experiments, except where otherwise stated. Iodine being volatile, it is to be expected that with hot solutions low results will be obtained.
Temperature15°20°40°60°80°"Hypo" required20.4 c.c.20.4 c.c.20.1 c.c.19.2 c.c.15.5 c.c.
These show that the temperature should not much exceed 20°.
Effect of Exposure of the Iodine Solution.—Twenty c.c. of the iodine were diluted to 100 c.c., and exposed for varying lengths of time in open beakers at the ordinary temperature, and then titrated.
Time exposed—1 day2 days3 days"Hypo" required20.4 c.c.16.1 c.c.13.6 c.c.9.4 c.c.
Effect of Varying Bulk.—These experiments were carried out in the usual way, bulk only varying.
Bulk100.0c.c.200.0c.c.300.0c.c.500.0c.c."Hypo" required20.4"20.4"20.4"20.4"
Effect of Varying Acid.—These experiments were under the usual conditions, the bulk being 100 c.c. The results were—
Acetic acid—1.5c.c.30.0c.c."Hypo" required20.4 c.c.20.7"20.7"Hydrochloric acid—1.5c.c.15.0c.c."Hypo" required20.4 c.c.20.6"20.9"Sulphuric acid—0.5c.c.20.0c.c."Hypo" required20.4 c.c.20.7"15.2"[86]Nitric acid—0.5c.c.10.0c.c."Hypo" required20.4 c.c.21.5"could not be titrated.
In the application of this titration to the assay of manganese ores, hydrochloric and hydriodic acids are the only ones likely to be present.
Effect of Alkalies.—On theoretical grounds the presence of these is known to be inadmissible. A solution rendered faintly alkaline with ammonia required only 11.2 c.c. of "hypo;" and another, with 0.5 gram of caustic soda, required 4.0 c.c. instead of 20.4 c.c. as in neutral solutions.
Effect of nearly Neutralising Hydrochloric Acid Solutions with Ammonia.—Provided care is taken not to addexcess of ammonia, this has a good effect, counteracting the interference of excess of acid. Thus 20 c.c. of iodine (as before) required 20.4 c.c. of "hypo;" with 15 c.c. of hydrochloric acid 20.7 c.c. were required, but with 15 c.c. of acid, nearly neutralised with dilute ammonia 20.4 c.c. were used.
Effect of the Addition of Starch.—The addition of varying quantities of starch has no effect, provided it is added when the titration is nearly finished, as the following experiments show:—
Starch added1.0c.c.5.0c.c.10.0c.c.50.0c.c."Hypo" required20.4"20.4"20.4"20.5"
But if the starch is added before the titration, the results are liable to error.
Starch added1.0c.c.50.0c.c."Hypo" required20.4"24.0"
The starch should be used fresh, and is best made on the day it is used; after four days the finishing point is not so good.
Effect of Varying Potassium Iodide.—An excess of iodide is always required to keep the iodine in solution; a larger excess has little effect.
Iodide added—1 gram20 grams"Hypo" required20.4 c.c.20.5 c.c.20.6 c.c.
The 20 c.c. of iodine used, itself contained 0.5 gram of potassium iodide.
Effect of Foreign Salts.—
Bicarbonate of soda added—0.5 gram1.5 gram5.0 grams"Hypo" required20.4 c.c.18.2 c.c.17.1 c.c.16.0 c.c.
The solution obviously must be free from bicarbonate of soda. This should be remembered, since when titrating arsenic assays with iodine it must be present; and students must avoid confounding the two titrations.
In some other experiments, in which 10 grams each of the salts were taken, the following results were obtained:—
Salt added—AmClAmNO3Am2SO4"Hypo" required20.4 c.c.20.5 c.c.20.3 c.c.20.2 c.c.
Salt addedNaClNaNO3Na2SO4"Hypo" required20.3 c.c.20.4 c.c.20.4 c.c.
Effect of Varying Iodine.—
Iodine added1.0c.c.10.0c.c.20.0c.c.50.0c.c.100.0c.c."Hypo" required1.3"10.2"20.4"51.0"102.0"
Determination of Dioxide in a Manganese Ore.—Weigh up 0.25 to 0.3 gram of the powdered ore; place in a flask, coverwith 10 c.c. of hydrochloric acid, and close the flask with a paraffined cork, and bulbs (as shown in fig. 60), having previously charged the bulb with 5 grams of potassium iodide in strong solution. Heat the flask, and boil cautiously for about fifteen minutes. Wash the contents of the bulbs into a large beaker, nearly (but not quite) neutralise with dilute ammonia, and titrate with the standard "hypo."
As an example, 0.2675 gram of pyrolusite was taken, and required 60.3 c.c. of standard "hypo" (100 c.c. equal 1.185 gram iodine, or 0.4042 gram MnO2), which equals 0.2437 gram of the dioxide or 91.1 per cent.
When compounds of manganese free from chlorides are boiled with nitric acid and dioxide of lead,[87]the manganese is converted into permanganic acid, which is soluble and tints the solution violet. The depth of colour depends on the amount of manganese present, and this should not much exceed 10 milligrams. A quantity of substance containing not more than this amount of manganese should be boiled for a few minutes with 25 c.c. of a solution containing 5 c.c. of nitric acid, and 10 or 20 c.c. of dilute sulphuric acid, with 2 or 3 grams of lead dioxide. Filter through asbestos, wash by decantation with dilute sulphuric acid, make up with distilled water[88]to a definite bulk, and take a measured portion for the colorimetric determination.
The standard solution of manganese is made by dissolving 0.1435 gram of permanganate of potash (KMnO4) in a little water acidulated with nitric acid, and diluting to 1 litre. One c.c. will contain 0.05 milligram of manganese.
1. What percentage of manganese (Mn) is contained in permanganate of potash (KMnO4)?
2. Ten c.c. of a solution of permanganate of potash is found to oxidise 10 c.c. of an acid solution of ferrous sulphate. The manganese is determined in the titrated solution by precipitation as dioxide and titrating. How much of the ferrous solution will be oxidised in the second titration?
3. What weight of potassium iodide would be just sufficient to absorb the chlorine evolved by 0.5 gram of pure dioxide of manganese?
4. What weight of iron must be dissolved up so as to have an excess of 0.25 gram after oxidation by 1 gram of pure dioxide?
5. What weight of the brown oxide, Mn2O4will be left on igniting 1 gram of the pure dioxide?
Chromium occurs in nature chiefly as chromite or chrome iron ore (FeO2Cr2O3, with more or less MgO and Al2O3), which is the chief ore. It is a constituent of some silicates, and is frequently met with in very small quantities in iron ores. It occurs as chromate in crocoisite (PbCrO4), and some other rare minerals.
The metal is used in steel-making. Steel containing about 0.5 per cent. of it is rendered very hard; but its chief value is in its salts, the chromates. These are highly-coloured compounds, generally red or yellow. Some of the insoluble chromates are used as pigments; chromate of lead or chrome-yellow is the most important. The soluble chromates, those of soda and potash, are valuable chemicals, and are largely used in the preparation of pigments, dyeing and tanning, and as oxidising agents.
Chromium forms two important classes of compounds—chromic salts, corresponding to the oxide Cr2O3, and chromates, which contain the trioxide CrO3. Solutions of chromic salts are green, whilst those of the chromates are yellow. Chromates are reduced to chromic salts by the action of most reducing agents in the presence of an acid; and this property is used in assaying for the volumetric determination of ferrous iron, &c. The chromates in solution are more stable than other similar oxidising agents, and consequently are generally used in the laboratory as one of the standard oxidising agents for volumetric analysis. They have the disadvantage of requiring an outside indicator. Bichromate of potash (K2Cr2O7) is the salt generally used for this purpose.
Chromic salts are oxidised to chromate by fusion with "fusion mixture" and nitre, or by treating with chlorine in an alkaline solution.
Chromic salts closely resemble those of ferric iron, and in the ordinary course of analysis chromic hydrate (green) is precipitated together with ferric hydrate, alumina, &c., on the addition of ammonic chloride and ammonia. The ignited oxide, Cr2O3, however, is not reduced on heating to redness in a current of hydrogen.
Detection.—Chromium is detected by fusing the powdered substance with "fusion mixture" and nitre. The melt is extracted with water and filtered. The filtrate is acidified with acetic acid, and treated with a few drops of a solution of leadacetate. A yellow precipitate indicates chromium. Substances containing chromium impart a green colour to the borax bead in both flames. Small quantities of chromate in neutral solution can be found by the dark or violet-red colouration imparted thereto on boiling with a dilute decoction of logwood.
Solution and Separation.—Chromates and chromic salts are generally soluble in water or dilute acids. Chrome iron ore, however, and ignited chromic oxide are insoluble; and the former presents considerable difficulty on attempting to open up by the usual methods. A large number of mixtures have been tried in order to get all the chromium in a soluble form. Among these are the following. One part of the very finely-powdered ore is fused with any of these mixtures.
(1) 10 parts of bisulphate of potash.(2) 5 parts of bisulphate of potash and 5 parts of potassium fluoride.(3) 5 parts of hydric potassic fluoride.(4) 12 parts of bisulphate of potash; and, afterwards, with 6 parts ofcarbonate of soda and 6 parts of nitre.(5) 8 parts of borax; afterwards, with carbonate of soda till it ceasesto effervesce; then, with 3 parts of carbonate of soda and 3 ofnitre.(6) 4 parts of borax and 6 parts of fusion mixture.(7) 12 parts of caustic potash.(8) 10 parts of caustic soda and 30 of magnesia.(9) 5 parts of caustic soda and 3 of magnesia.(10) 2 parts of carbonate of soda and 1 of lime.(11) 6 parts of soda-lime and 2 of chlorate of potash.(12) Sodium peroxide.
Of these, numbers 1, 2, and 3 yield the chromium in a form soluble in dilute acids, as chromic salt. The rest in a form soluble in water, as potassium or sodium chromate.
On boiling an insoluble chromium compound with chlorate of potash and nitric acid, the chromium passes into solution as chromate. This method, however, does not answer for chrome iron ore. In the fusion methods the ore must be very finely powdered, well mixed with the fluxes, and subjected to a prolonged fusion in a platinum vessel at a high temperature. Undecomposed particles require re-fusion.
The aqueous extract containing the chromate is ready for volumetric work, except in those cases where nitre has been used. For gravimetric work the solution is acidified with hydrochloric acid, then mixed with ammonia in slight excess, boiled, and filtered. The filtrate is acidified with hydrochloric acid, and treated with sulphuretted hydrogen, warmed, rendered slightly alkaline with ammonia, and the gas again passed. The chromium is precipitated as chromic hydrate mixed with sulphur from thereduction with sulphuretted hydrogen. It is filtered off, washed with hot water, and ignited. It is weighed as chromic oxide.
The solution containing the chromium, freed from other metals and earths and in the form of (green) chromic salt, is heated to boiling. If any chromate is present reduce it with sodium sulphite or sulphuretted hydrogen. Add ammonia in slight excess, boil till the liquid is free from a red tint, and allow to settle for a few minutes. Filter, wash with hot water, dry, and ignite strongly in a loosely-covered crucible. Cool, and weigh. The substance is chromic oxide, Cr2O3, and contains 68.62 per cent. of chromium. It is a dark-green powder insoluble in acids.
When, as is generally the case, the chromium exists altogether as chromate (phosphates and arsenates being absent) it is best to proceed as follows:—Render the solution acid with acetic acid, then add sodium acetate to the solution and heat nearly to boiling; next treat with a slight excess of acetate of lead, and boil. Allow to settle, and filter. Wash the precipitate with hot water, dry in the water-oven or at a low temperature. Transfer the precipitate to a weighed Berlin crucible, burn the filter separately, ignite below redness, cool in the desiccator, and weigh. The substance is lead chromate, PbCrO4, and contains 16.1 per cent. of chromium, or 23.53 per cent. of chromic oxide (Cr2O3).
This is based on the oxidation of ferrous iron by the solution containing the chromium as chromate. A known weight of iron (0.5, 1, or 1.5 gram, according to the quantity of chromate) is dissolved in 50 c.c. of dilute sulphuric acid. The solution containing the chromate is added, and the remaining ferrous iron titrated with the permanganate or bichromate of potassium solution, as described underIron. The iron thus found is deducted from that taken, and the difference gives the iron oxidised by the chromate. This multiplied by 0.3101 gives the chromium, Cr, and when multiplied by 0.4529 gives the chromic oxide, Cr2O3.
Small quantities of chromium may be determined, after conversion into chromate, colorimetrically. The solution, which should not contain more than a few milligrams in 100 c.c., is acidified with acetic acid and compared against an equal volumeof water rendered acid with acetic acid and tinted with a standard bichromate of potassium solution. This standard bichromate is made by dissolving 2.827 grams of the salt in water and diluting to 1 litre. One c.c. will contain 1 milligram of chromium, Cr. The manner of working this assay is the same as that adopted in the other colorimetric processes.
Determination of Chromium in Steel.[89]—Weigh up 2.4 grams, dissolve in hydrochloric acid, and evaporate to dryness. Fuse with sodium carbonate and nitre, extract with water, and make up to 301 c.c. Take 250 c.c. of the clear liquor, boil with hydrochloric acid, add sodium phosphate, and then ammonia in slight excess. Heat till clear. Filter off the precipitate, dissolve it in hydrochloric acid, and evaporate to dryness. Take up with a little acid, filter, and precipitate with a slight excess of ammonia. Wash, ignite, and weigh as chromium phosphate (3Cr2O3,2P2O5), which contains 42.2 per cent. of chromium.
Vanadium occurs in certain rare minerals, such as vanadinite (3Pb3(VO4)2.PbCl2), a vanadate of lead; mottramite, a vanadate of copper and lead; and dechenite, a vanadate of lead and zinc. It is occasionally found in iron and copper ores and in the slags from them. In Spanish copper-precipitates it is found along with chromium, and is probably derived from the iron used for precipitating. The vanadates, like the chromates, are coloured compounds, generally yellow or red. On reduction, blue solutions are got. In their general reactions vanadates resemble phosphates.
Vanadium is detected by the red colouration produced by passing sulphuretted hydrogen into ammoniacal solutions for some time. On adding an acid to the filtered solution a brown precipitate of the sulphide is produced. This gives with borax a colourless bead in the oxidising, and a green one in the reducing, flame.
It is separated by fusing the ore with potassic nitrate, extracting with water and precipitating with baric chloride. The precipitate is boiled with dilute sulphuric acid, filtered, neutralised with ammonia, and saturated with ammonic chloride. Ammonium vanadate separates out. It is filtered off, ignited, and weighed as vanadic oxide, V2O5, containing 56.18 per cent. of vanadium.
Molybdenum occurs in nature chiefly as molybdenite (MoS2); it also occurs in wulfenite, a molybdate of lead (PbMoO4), and in molybdic ochre (MoO3).
Molybdate of ammonia is an important reagent in the determination of phosphates, the manufacture of which compound is the chief purpose to which molybdenum is applied.
Iron and copper ores frequently contain molybdenum, sometimes in quantity; consequently it is met with in slags and pig-iron.
Molybdenum forms several series of salts. In those corresponding to the lower oxides it is basic; but the trioxide (MoO3) is the acid oxide which forms a series of salts called the molybdates. All molybdenum compounds are converted into the trioxide by boiling with nitric acid. The trioxide is a white powder readily dissolved by ammonia. It fuses at a red heat, and volatilises freely in contact with air. It is slightly soluble in water.
Molybdates are easily reduced, with the production of coloured solutions, by most reducing agents. Sulphuretted hydrogen first produces a blue tint, and then precipitates a brown sulphide. The precipitation as sulphide is only complete on prolonged treatment; a green colour indicates that some molybdenum still remains in solution. The precipitated sulphide is soluble in ammonium sulphide.
Detection.—Molybdenum is detected by its behaviour with sulphuretted hydrogen. Molybdenite can only be mistaken for graphite, from which it is easily distinguished by yielding sulphur dioxide on roasting, and by giving, on charcoal, a yellowish white incrustation, which becomes blue on touching it for a moment with the reducing flame. The borax-bead is colourless in the oxidising, and dark-brown in the reducing, flame.
The solution containing the molybdate is neutralised and treated with an excess of mercurous nitrate. The precipitate is allowed to settle for some time, filtered, and washed with a dilute solution of mercurous nitrate. Then it is dried, transferred to a weighed Berlin crucible containing ignited oxide of lead, mixed, ignited, and weighed. The increase in weight gives the amount of trioxide, MoO3. This contains 66.7 per cent. of molybdenum.
Uranium occurs chiefly as pitchblende, which is an impure oxide (U3O8). It is also found as sulphate in uranochre, johannite, &c.; and as phosphate in the uranites, torbernite (hydrated phosphate of uranium and copper), and autunite (hydrated phosphate of uranium and lime). It also occurs in some rarer minerals.
The oxide is used for colouring glass; and the nitrate and acetate are used as reagents. "Uranium yellow," used for enamel painting, is sodium uranate. The uranates, in which the oxide of uranium acts as an acid, are mostly insoluble and of secondary importance.
Uranium forms two families of salts, uranous and uranic; corresponding to the oxides UO2and UO3respectively. The former are generally green and the latter yellow. Uranous salts are converted into uranic by boiling with nitric acid or other oxidising agents. Uranic salts, on the other hand, are easily reduced by sulphuretted hydrogen, stannous chloride or zinc. This property is made use of in determining the quantity of uranium in pure solutions by titrating with permanganate of potassium solution as in the case with iron.
Detection.—The most characteristic reaction of the uranium compounds is their behaviour in the presence of alkaline carbonates in which they are freely soluble; even ammonium sulphide will not precipitate uranium from these solutions. On neutralising the carbonate with an acid a uranate of the alkali is precipitated. Ammonia or sodic hydrate (free from carbonates) give yellow precipitates, which are insoluble in excess of the reagent, but are soluble in acids. Ferrocyanide of potassium gives a reddish-brown precipitate. Uranium colours the borax-bead yellowish-green in the oxidising, and green in the reducing, flame.
Solution and Separation.—The compounds of uranium are soluble in acids. Powder the substance and evaporate with an excess of nitric acid. Take up with hydrochloric acid, dilute, pass sulphuretted hydrogen, and filter. Peroxidise the filtrate with a little nitric acid, add an excess of ammonic carbonate and some ammonium sulphide, and filter. Render the solution acid, boil; and precipitate the uranium by means of ammonia. Filter off, and wash it with dilute ammonic chloride. Ignite, and weighas protosesqui-oxide, U3O8.
The solution containing the uranium free from other metals is, if required, first peroxidised by boiling with nitric acid. Ammonia in slight excess is added to the nearly-boiling solution. A yellow precipitate is formed, which is filtered off hot and washed with a dilute solution of ammonium chloride. The precipitate is dried and ignited; and weighed as U3O8, which contains 84.8 per cent. of uranium.
This is based on the precipitation of uranium as phosphate from acetic acid solutions and the recognition of complete precipitation by testing with potassic ferrocyanide; it is the converse of the process for the volumetric determination of phosphate.
The standard solution of phosphateis prepared by dissolving 29.835 grams of hydric sodic phosphate (Na2HPO4.12H2O) in water and diluting to 1 litre. 100 c.c. will be equivalent to 2 grams of uranium.
Take 1 gram of the sample (or, if poor in uranium, 2 grams) and separate the uranium as described. Dissolve the precipitate in nitric acid and evaporate to a small bulk, add 2 grams of sodium acetate, dilute with water to 100 c.c., and boil. Titrate the boiling solution with the sodium phosphate till it ceases to give a brown colouration with potassium ferrocyanide. Calculate the percentage in the usual way.