Vogel: Textbook of Inorganic Analysis: 4^th Ed. P 348-351,370-372

 

B.4 OXIDATION - REDUCTION TITRATIONS

 

In the following Sections we are concerned with the titration of reducing agents with oxidising agents such as potassium permanganate, potassium dichromate, cerium(IV) sulphate, iodine, potassium iodate and potassium bromate, and with the titration of oxidising agents by reducing reagents such as arsenic(III) oxide and sodium thiosulphate.

 

The relevant theoretical Sections (X, 31-33) should be studied, and it should also be noted that in many cases, before titration with an oxidising reagent is carried out, it is necessary to ensure that the substance to be titrated is in a suitable lower oxidation state, i.e., it may be necessary to reduce the test solution before titration can be carried out. A selection of methods for carrying out such reductions is given at the end of this Chapter (Sections X, 142-6).

 

Oxidations with potassium permanganate

 

X, 90. DISCUSSION. This valuable and powerful oxidising agent was first introduced into titrimetric analysis by F. Margueritte for the titration of iron(II). In acid solutions, the reduction can be represented by the following equation:

 

MnO₄^- + 8H⁺ + 5e  « Mn²⁺ +4H₂O

 

from which it follows that the equivalent is one-fifth of the mole, i.e. 158.03/5, or 31.606. The standard potential in acid solution, E°, has been calculated to be 1.51 volts, hence the permanganate ion in acid solution is a strong oxidising agent.

 

Sulphuric acid is the most suitable acid, as it has no action upon permanganate in dilute solution. With hydrochloric acid, there is the likelihood of the reaction:

 

2MnO₄^- + 10Cl^- + 16H⁺ = 2Mn²⁺ + 5C1₂ + 8H₂O

 

taking place, and some permanganate may be consumed in the formation of chlorine. This reaction is particularly liable to occur with iron salts unless special precautions are adopted (see below). With a small excess of free acid, a very dilute solution, low temperature and slow titration with constant shaking, the danger from this cause is minimised. There are, however, some titrations, such as those with arsenic(III) oxide, trivalent antimony, and hydrogen peroxide, which can be carried out in the presence of hydrochloric acid.

 

In the analysis of iron ores, solution is frequently effected in concentrated hydrochloric acid; the iron(III) is reduced and the iron(II) is then determined in the resultant solution. To do this, it is best to add about 25 cm³ of Zimmermann and Reinhardt's solution (this is sometimes termed preventive solution), which is prepared by dissolving 50 g crystallised manganese(II) sulphate MnSO₄,4H₂O in 250 cm³ water, adding a cooled mixture of 100 cm³ concentrated sulphuric acid and 300 cm³ water, followed by 100 cm³ syrupy phosphoric acid. The manganese(II) sulphate lowers the reduction potential of the MnO₄^--Mn(II) couple (compare Sections II, 23-24) and thereby makes it a weaker oxidising agent; the tendency of the permanganate ion to oxidise chloride ion is thus reduced. It has been stated that a further function of the manganese(II) sulphate is to supply an adequate concentration of Mn²⁺ ions to react with any local excess of permanganate ion. Mn(III) is probably formed in the reduction of permanganate ion to manganese(II); the Mn(II), and also the phosphoric acid, exert a depressant effect upon the potential of the Mn(III)-Mn(II) couple, so that Mn(III) is reduced by Fe²⁺ ion rather than by chloride ion. The phosphoric acid combines with the yellow Fe²⁺ ion to form the complex ion [Fe(HPO₄)]⁺, thus rendering the end-point more clearly visible. The phosphoric acid lowers the reduction potential of the Fe(III)-Fe(II) system by complexation, and thus tends to increase the reducing power of the Fe²⁺ ion. Under these conditions permanganate ion oxidises iron(II) rapidly and reacts only slowly with chloride ion.

 

For the titration of colourless or slightly coloured solutions, the use of an indicator is unnecessary, since as little as 0.01 cm³ of 0.01N-potassium permanganate imparts a pale-pink colour to 100 cm³ of water. The intensity of the colour in dilute solutions may be enhanced, if desired, by the addition of a redox indicator (such as sodium diphenylamine sulphonate, N-phenylanthranilic acid, or ferroin) just before the end-point of the reaction; this is usually not required, but is advantageous if more dilute solutions of permanganate are used.

 

Potassium permanganate also finds some application in strongly alkaline solutions. Here two consecutive partial reactions take place:

(i) the relatively rapid reaction:

 

MnO₄^- + e « MnO₄^2-

 

and (ii) the relatively slow reaction:

 

MnO₄^2- +2H₂O + 2e « MnO₂+4OH^-

 

The standard potential Eє of reaction (i) is 0.56 volt and of reaction (ii) 0.60 volt. By suitably controlling the experimental conditions (e.g., by the addition of barium ions, which form the sparingly soluble barium manganate as a fine, granular precipitate), reaction (i) occurs almost exclusively; the equivalent is then 1 mole. In moderately alkaline solutions permanganate is reduced quantitatively to manganese dioxide. The half-cell reaction is:

 

MnO₄^- +2H₂O + 3e « MnO₂ +4OH^-

 

and the standard potential Eє is 0.59 volt.

 

Potassium permanganate is not a primary standard. It is difficult to obtain the substance perfectly pure and completely free from manganese dioxide. Moreover, ordinary distilled water is likely to contain reducing substances (traces of organic matter, etc.) which will react with the potassium permanganate to form manganese dioxide. The presence of the latter is very objectionable because it catalyses the auto-decomposition of the permanganate solution on standing. The decomposition:

 

4MnO₄^- + 2H₂O  =  4MnO₂+3O₂+4OH^-

 

is catalysed by solid manganese dioxide. Permanganate is inherently unstable in the presence of manganese(II) ions:

 

2MnO₄^-+3Mn²⁺ +2H₂O = 5MnO₂ + 4H⁺;

 

this reaction is slow in acid solution, but is very rapid in neutral solution. For these reasons, potassium permanganate solution is rarely made up by dissolving weighed amounts of the highly purified (e.g., A.R.) solid in water; it is more usual to heat a freshly prepared solution to boiling and keep it on the steam bath for an hour or so, and then filter the solution through a non-reducing filtering medium, such as purified glass wool or a sintered glass filtering crucible (porosity No. 4).

 

 Alternatively, the solution may be allowed to stand for 2-3 days at room temperature before filtration. The glass-stoppered bottle or flask should bc carefully freed from grease and prior deposits of manganese dioxide: this may be done by rinsing with dichromate-sulphuric acid cleaning mixture and then thoroughly with distilled water. Acidic and alkaline solutions are less stable than neutral ones. Solutions of permanganate should be protected from unnecessary exposure to light; a dark-coloured bottle is recommended. Diffuse daylight causes no appreciable decomposition, but bright sunlight slowly decomposes even pure solutions.

 

Potassium permanganate solutions may be standardised using arsenic(III) oxide or sodium oxalate as primary standards: secondary standards include metallic iron, and iron(II) ethylenediammonium sulphate (or ethylenediamine iron(II) sulphate), FeSO₄,C₂H₄(NH₃)₂SO₄, 4H₂O.

 

Of these substances sodium oxalate was formerly regarded as the most trustworthy, since it is readily obtained pure and anhydrous, and the ordinary A.R. substance has a purity of at least 99.9 per cent. The experimental procedure hitherto employed was due to R. S. McBride. A solution of the oxalate, acidified with dilute sulphuric acid and warmed to 80-90 єC, was titrated with the permanganate solution slowly (10-15 cm³ per minute) and with constant stirring until the first permanent faint pink colour was obtained; the temperature near the end-point was not allowed to fall below 60 єC. R. M. Fowler and H. A. Bright have, however, shown that with McBride's procedure the results may be 0.1 - 0.45 per cent high; the titre depends upon the acidity, the temperature, the rate of addition of the permanganate solution, and upon the speed of stirring. These authors recommend a more rapid addition of 90-95 per cent of the permanganate solution (about 25-35 cm³ per minute) to a solution of sodium oxalate in M­-sulphuric acid at 25-30 єC, the solution is then warmed to 55-60 єC and the titration completed, the last 0.5- 1 cm³ portion being added dropwise. The method is accurate to 0.06 per cent. Full experimental details are given in Procedure B below.

 

2Na⁺ +C₂O₄^2- +2H⁺ ® H₂C₂O₄ + 2Na⁺

 

2MnO₄^- + 5H₂C₂O₄ + 6H⁺® 2Mn²⁺ + 10CO₂ + 8H₂O

 

It should be mentioned that if oxalate is to be determined it is often not convenient to use the room-temperature technique for unknown amounts of oxalate. The permanganate solution may then be standardised against sodium oxalate at about 80 єC using the same procedure in the standardisation as in the analysis.

 

The procedure of H. A. Bright, which utilises arsenic(III) oxide as a primary standard and potassium iodide or potassium iodate as a catalyst for the reaction, is more convenient in practice and is a trustworthy method for the standardisation of permanganate solutions. A.R. arsenic(III) oxide has a purity of at least 99.8 per cent, and the results by this method agree to within 1 part in 3000 with the sodium oxalate procedure of Fowler and Bright. Full experimental details are given in Procedure A (Section X, 92).

 

As₂O₃+4OH^- ® 2HAsO₃^2- + H₂O

 

5H₃AsO₃ + 2MnO₄^- + 6H⁺ = 5H₃AsO₄ + 2Mn²⁺ + 3H₂O

 

Potassium iodide, if specially purified, may be used as a primary standard. For many practical purposes, the dry A.R. reagent is sufficiently pure. The potentiometric method (see Chapter XIV) should be employed: a bright platinum indicator electrode and a saturated calomel electrode are required. The concentration of the sulphuric acid should be about O.4M.

 

10I^- + 2MnO₄^- + 16H⁺ ® 5I₂ + 2Mn²⁺ + 8H₂O

 

Iron wire of 99.9 per cent purity is available commercially and the A.R. reagent is a suitable standard, particularly if the potassium permanganate solution is subsequently to be employed in the determination of iron. If the wire exhibits any sign of rust, it should be drawn between two pieces of fine emery cloth, and then wiped with a clean, dry cloth before use. The reaction which occurs is:

 

MnO₄^- + 5Fe²⁺ +8H⁺ ® Mn²⁺ + 5Fe³⁺ + 4H₂O

 

Ethylenediammonium iron(II) sulphate, FeSO₄, C₂H₄(NH₃)₂SO₄, 4H₂O is relatively stable and has a high molecular weight (382.16). The preparation is as follows.

 

To 10.0 g of a 99 per cent solution of ethylenediamine, add 60 cm³ 6N-sulphuric acid and 46.3 g A.R. iron(II) sulphate heptahydrate. Dilute to 300 cm³ with distilled water, and to the resulting solution introduce 300 cm³ of ethanol slowly and with constant stirring. Filter through a sintered glass funnel, wash the precipitate with 50 per cent ethanol, and redissolve it in slightly acidulated water. Add two-thirds the volume of ethanol. Filter again as before, and wash the solid successively with 65 per cent ethanol and 95 per cent ethanol. Dry in the air or at 50 єC for about 12 hours. The yield is about 50 g.

 

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