!Method A!
The procedures, as here prescribed, are applicable to iron ores in general, provided these ores contain no const.i.tuents which are reduced by zinc or stannous chloride and reoxidized by permanganates. Many iron ores contain t.i.tanium, and this element among others does interfere with the determination of iron by the process described.
If, however, the solutions of such ores are treated with sulphureted hydrogen or sulphurous acid, instead of zinc or stannous chloride to reduce the iron, and the excess reducing agent removed by boiling, an accurate determination of the iron can be made.
PROCEDURE.--Grind the mineral to a fine powder. Weigh out two portions of about 0.5 gram each into small porcelain crucibles. Roast the ore at dull redness for ten minutes (Note 1), allow the crucibles to cool, and place them and their contents in ca.s.seroles containing 30 cc. of dilute hydrochloric acid (sp. gr. 1.12).
Proceed with the solution of the ore, and the treatment of the residue, if necessary, exactly as described for the bichromate process on page 56. When solution is complete, add 6 cc. of concentrated sulphuric acid to each ca.s.serole, and evaporate on the steam bath until the solution is nearly colorless (Note 2). Cover the ca.s.seroles and heat over the flame of the burner, holding the ca.s.serole in the hand and rotating it slowly to hasten evaporation and prevent spattering, until the heavy white fumes of sulphuric anhydride are freely evolved (Note 3). Cool the ca.s.seroles, add 100 cc. of water (measured), and boil gently until the ferric sulphate is dissolved; pour the warm solution through the reductor which has been previously washed; proceed as described under standardization, taking pains to use the same volume and strength of acid and the same volume of wash-water as there prescribed, and t.i.trate with the permanganate solution in the reductor flask, using the ferrous sulphate solution if the end-point should be overstepped.
From the corrected volume of permanganate solution used, calculate the percentage of iron (Fe) in the limonite.
[Note 1: The preliminary roasting is usually necessary because, even though the sulphuric acid would subsequently char the carbonaceous matter, certain nitrogenous bodies are not thereby rendered insoluble in the acid, and would be oxidized by the permanganate.]
[Note 2: The temperature of the steam bath is not sufficient to volatilize sulphuric acid. Solutions may, therefore, be left to evaporate overnight without danger of evaporation to dryness.]
[Note 3: The hydrochloric acid, both free and combined, is displaced by the less volatile sulphuric acid at its boiling point. Ferric sulphate separates at this point, since there is no water to hold it in solution and care is required to prevent b.u.mping. The ferric sulphate usually has a silky appearance and is easily distinguished from the flocculent silica which often remains undissolved.]
!Zimmermann-Reinhardt Procedure!
!Method (B)!
PROCEDURE.--Grind the mineral to a fine powder. Weigh out two portions of about 0.5 gram each into small porcelain crucibles. Proceed with the solution of the ore, treat the residue, if necessary, and reduce the iron by the addition of stannous chloride, followed by mercuric chloride, as described for the bichromate process on page 56. Dilute the solution to about 400 cc. with cold water, add 10 cc. of the manganous sulphate t.i.trating solution (Note 1, page 68) and t.i.trate with the standard pota.s.sium permanganate solution to a faint pink (Note 1).
From the standardization data already obtained calculate the percentage of iron (Fe) in the limonite.
[Note 1: It has already been noted that hydrochloric acid reacts slowly in cold solutions with pota.s.sium permanganate. It is, however, possible to obtain a satisfactory, although somewhat fugitive end-point in the presence of manganous sulphate and phosphoric acid.
The explanation of the part played by these reagents is somewhat obscure as yet. It is possible that an intermediate manganic compound is formed which reacts rapidly with the ferrous compounds--thus in effect catalyzing the oxidizing process.
While an excess of hydrochloric acid is necessary for the successful reduction of the iron by stannous chloride, too large an amount should be avoided in order to lessen the chance of reduction of the permanganate by the acid during t.i.tration.]
DETERMINATION OF THE OXIDIZING POWER OF PYROLUSITE
INDIRECT OXIDATION
Pyrolusite, when pure, consists of manganese dioxide. Its value as an oxidizing agent, and for the production of chlorine, depends upon the percentage of MnO_{2} in the sample. This percentage is determined by an indirect method, in which the manganese dioxide is reduced and dissolved by an excess of ferrous sulphate or oxalic acid in the presence of sulphuric acid, and the unused excess determined by t.i.tration with standard permanganate solution.
PROCEDURE.--Grind the mineral in an agate mortar until no grit whatever can be detected under the pestle (Note 1). Transfer it to a stoppered weighing-tube, and weigh out two portions of about 0.5 gram into beakers (400-500 cc.) Read Note 2, and then calculate in each case the weight of oxalic acid (H_{2}C_{2}O_{4}.2H_{2}O) required to react with the weights of pyrolusite taken. The reaction involved is
MnO_{2} + H_{2}C_{2}O_{4}(2H_{2}O) + H_{2}SO_{4} --> MnSO_{4} + 2CO_{2} + 4H_{2}O.
Weigh out about 0.2 gram in excess of this quant.i.ty of !pure! oxalic acid into the corresponding beakers, weighing the acid accurately and recording the weight in the notebook. Pour into each beaker 25 cc. of water and 50 cc. of dilute sulphuric acid (1:5), cover and warm the beaker and its contents gently until the evolution of carbon dioxide ceases (Note 3). If a residue remains which is sufficiently colored to obscure the end-reaction of the permanganate, it must be removed by filtration.
Finally, dilute the solution to 200-300 cc., heat the solution to a temperature just below boiling, add 15 cc. of a manganese sulphate solution and while hot, t.i.trate for the excess of the oxalic acid with standard permanganate solution (Notes 4 and 5).
From the corrected volume of the solution required, calculate the amount of oxalic acid undecomposed by the pyrolusite; subtract this from the total quant.i.ty of acid used, and calculate the weight of manganese dioxide which would react with the balance of the acid, and from this the percentage in the sample.
[Note 1: The success of the a.n.a.lysis is largely dependent upon the fineness of the powdered mineral. If properly ground, solution should be complete in fifteen minutes or less.]
[Note 2: A moderate excess of oxalic acid above that required to react with the pyrolusite is necessary to promote solution; otherwise the residual quant.i.ty of oxalic acid would be so small that the last particles of the mineral would scarcely dissolve. It is also desirable that a sufficient excess of the acid should be present to react with a considerable volume of the permanganate solution during the t.i.tration, thus increasing the accuracy of the process. On the other hand, the excess of oxalic acid should not be so large as to react with more of the permanganate solution than is contained in a 50 cc. burette. If the pyrolusite under examination is known to be of high grade, say 80 per cent pure, or above the calculation of the oxalic acid needed may be based upon an a.s.sumption that the mineral is all MnO_{2}. If the quality of the mineral is unknown, it is better to weigh out three portions instead of two and to add to one of these the amount of oxalic prescribed, a.s.suming complete purity of the mineral. Then run in the permanganate solution from a pipette or burette to determine roughly the amount required. If the volume exceeds the contents of a burette, the amount of oxalic acid added to the other two portions is reduced accordingly.]
[Note 3: Care should be taken that the sides of the beaker are not overheated, as oxalic acid would be decomposed by heat alone if crystallization should occur on the sides of the vessel. Strong sulphuric acid also decomposes the oxalic acid. The dilute acid should, therefore, be prepared before it is poured into the beaker.]
[Note 4: Ferrous ammonium sulphate, ferrous sulphate, or iron wire may be subst.i.tuted for the oxalic acid. The reaction is then the following:
2 FeSO_{4} + MnO_{2} + 2H_{2}SO_{4} --> Fe_{2}(SO_{4})_{3} + 2H_{2}O
The excess of ferrous iron may also be determined by t.i.tration with pota.s.sium bichromate, if desired. Care is required to prevent the oxidation of the iron by the air, if ferrous salts are employed.]
[Note 5: The oxidizing power of pyrolusite may be determined by other volumetric processes, one of which is outlined in the following reactions:
MnO_{2} + 4HCl --> MnCl_{2} + Cl_{2} + 2H_{2}O Cl_{2} + 2KI --> I_{2} + 2KCl I_{2} + 2Na_{2}S_{2}O_{3} --> Na_{2}S_{4}O_{6} + 2NaI.
The chlorine generated by the pyrolusite is pa.s.sed into a solution of pota.s.sium iodide. The liberated iodine is then determined by t.i.tration with sodium thiosulphate, as described on page 78. This is a direct process, although it involves three steps.]
IODIMETRY
The t.i.tration of iodine against sodium thiosulphate, with starch as an indicator, may perhaps be regarded as the most accurate of volumetric processes. The thiosulphate solution may be used in both acid and neutral solutions to measure free iodine and the latter may, in turn, serve as a measure of any substance capable of liberating iodine from pota.s.sium iodide under suitable conditions for t.i.tration, as, for example, in the process outlined in Note 5 on page 74.
The fundamental reaction upon which iodometric processes are based is the following:
I_{2} + 2 Na_{2}S_{2}O_{3} --> 2 NaI + Na_{2}S_{4}O_{6}.
This reaction between iodine and sodium thiosulphate, resulting in the formation of the compound Na_{2}S_{4}O_{6}, called sodium tetrathionate, is quant.i.tatively exact, and differs in that respect from the action of chlorine or bromine, which oxidize the thiosulphate, but not quant.i.tatively.
NORMAL SOLUTIONS OF IODINE AND SODIUM THIOSULPHATE
If the formulas of sodium thiosulphate and sodium tetrathionate are written in a manner to show the atoms of oxygen a.s.sociated with sulphur atoms in each, thus, 2(Na_{2}).S_{2}O_{2} and Na_{2}O.S_{4}O_{5}, it is plain that in the tetrathionate there are five atoms of oxygen a.s.sociated with sulphur, instead of the four in the two molecules of the thiosulphate taken together. Although, therefore, the iodine contains no oxygen, the two atoms of iodine have, in effect, brought about the addition of one oxygen atoms to the sulphur atoms. That is the same thing as saying that 253.84 grams of iodine (I_{2}) are equivalent to 16 grams of oxygen; hence, since 8 grams of oxygen is the basis of normal solutions, 253.84/2 or 126.97 grams of iodine should be contained in one liter of normal iodine solution. By a similar course of reasoning the conclusion is reached that the normal solution of sodium thiosulphate should contain, per liter, its molecular weight in grams. As the thiosulphate in crystalline form has the formula Na_{2}S_{2}O_{3}.5H_{2}O, this weight is 248.12 grams. Tenth-normal or hundredth-normal solutions are generally used.
PREPARATION OF STANDARD SOLUTIONS
!Approximate Strength, 0.1 N!
PROCEDURE.--Weigh out on the rough balances 13 grams of commercial iodine. Place it in a mortar with 18 grams of pota.s.sium iodide and triturate with small portions of water until all is dissolved. Dilute the solution to 1000 cc. and transfer to a liter bottle and mix thoroughly (Note 1).[1]
[Footnote 1: It will be found more economical to have a considerable quant.i.ty of the solution prepared by a laboratory attendant, and to have all unused solutions returned to the common stock.]
Weigh out 25 grams of sodium thiosulphate, dissolve it in water which has been previously boiled and cooled, and dilute to 1000 cc., also with boiled water. Transfer the solution to a liter bottle and mix thoroughly (Note 2).