Chemical elements
  Iron
    History of Iron
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    Iron Salts
    PDB 101m-1aeb
    PDB 1aed-1awd
    PDB 1awp-1beq
    PDB 1bes-1c53
    PDB 1c6o-1ci6
    PDB 1cie-1cry
    PDB 1csu-1dfx
    PDB 1dgb-1dry
    PDB 1ds1-1e08
    PDB 1e0z-1ehj
    PDB 1ehk-1f5o
    PDB 1f5p-1fnp
    PDB 1fnq-1fzi
    PDB 1g08-1gnl
    PDB 1gnt-1h43
    PDB 1h44-1hdb
    PDB 1hds-1i5u
    PDB 1i6d-1iwh
    PDB 1iwi-1jgx
    PDB 1jgy-1k2o
    PDB 1k2r-1kw6
    PDB 1kw8-1lj0
    PDB 1lj1-1m2m
    PDB 1m34-1mko
    PDB 1mkq-1mun
    PDB 1muy-1n9x
    PDB 1naz-1nx4
    PDB 1nx7-1ofe
    PDB 1off-1p3t
    PDB 1p3u-1pmb
    PDB 1po3-1qmq
    PDB 1qn0-1ra0
    PDB 1ra5-1rxg
    PDB 1ry5-1smi
    PDB 1smj-1t71
    PDB 1t85-1u8v
    PDB 1u9m-1uyu
    PDB 1uzr-1vxf
    PDB 1vxg-1wri
    PDB 1wtf-1xlq
    PDB 1xm8-1y4r
    PDB 1y4t-1ygd
    PDB 1yge-1z01
    PDB 1z02-2a9e
    PDB 2aa1-2azq
    PDB 2b0z-2boz
    PDB 2bpb-2ca3
    PDB 2ca4-2cz7
    PDB 2czs-2dyr
    PDB 2dys-2ewk
    PDB 2ewu-2fwl
    PDB 2fwt-2gl3
    PDB 2gln-2hhb
    PDB 2hhd-2ibn
    PDB 2ibz-2jb8
    PDB 2jbl-2mgh
    PDB 2mgi-2o01
    PDB 2o08-2ozy
    PDB 2p0b-2q0i
    PDB 2q0j-2r1h
    PDB 2r1k-2spm
    PDB 2spn-2vbd
    PDB 2vbp-2vzb
    PDB 2vzm-2wiv
    PDB 2wiy-2xj5
    PDB 2xj6-2ylj
    PDB 2yrs-2zon
    PDB 2zoo-3a17
    PDB 3a18-3aes
    PDB 3aet-3bnd
    PDB 3bne-3cir
    PDB 3ciu-3dax
    PDB 3dbg-3e1p
    PDB 3e1q-3eh4
    PDB 3eh5-3fll
    PDB 3fm1-3gas
    PDB 3gb4-3h57
    PDB 3h58-3hrw
    PDB 3hsn-3ir6
    PDB 3ir7-3k9y
    PDB 3k9z-3l4p
    PDB 3l61-3lxi
    PDB 3lyq-3mm8
    PDB 3mm9-3n62
    PDB 3n63-3nlo
    PDB 3nlp-3o0f
    PDB 3o0r-3p6o
    PDB 3p6p-3prq
    PDB 3prr-3sel
    PDB 3sik-3una
    PDB 3unc-4blc
    PDB 4cat-4erg
    PDB 4erm-4nse
    PDB 4pah-8cat
    PDB 8cpp-9nse

Estimation of Iron





Estimation of Iron by Gravimetric Methods

If the iron is already in solution, it is first oxidised to the ferric condition and precipitated as ferric hydroxide by addition of ammonia to the boiling solution. The precipitate is well washed, dried in an oven, ignited in a crucible, and weighed as anhydrous ferric oxide, Fe2O3.

If the iron is not already in solution, the solid to be analysed is digested with hydrochloric acid, or aqua regia, or is brought into a soluble condition by fusion with potassium carbonate or hydrogen sulphate. Silica is filtered off, any copper, lead, etc., precipitated with hydrogen sulphide, and the iron oxidised to the ferric condition with nitric acid. Addition of ammonium chloride and hydroxide precipitates the iron, which is filtered off and weighed as Fe2O3. If aluminium and chromium are likely to be present, these are first removed, as in qualitative analysis, by adding sodium peroxide to the precipitated hydroxide, prior to ignition. After washing away any sodium aluminate and chromate, the pure ferric hydroxide is ignited, as already indicated.

If manganese was originally present, some of it will precipitate out with the iron, and be weighed as Mn3O4 along with the Fe2O3. If it is desired to remove the manganese, the hydroxides, prior to ignition, are dissolved in a minimum quantity of hydrochloric acid, and ammonium carbonate added, under constant agitation, until the precipitate first formed just re-dissolves, leaving the liquid slightly opalescent. Acetic acid and ammonium acetate are now added, the solution boiled for a few moments, filtered hot, and the precipitate, consisting of basic ferric acetate, washed twice with boiling water. The filtrate will be colourless if the operations have been successfully carried out, and contains the bulk of the manganese as acetate. A small quantity of manganese, however, will still be entangled in the precipitate, which latter is therefore re-dissolved in acid and precipitated a second time in a precisely- similar manner, washed, and ignited to Fe2O3.

" Cupferron " or amino nitroso phenyl hydroxylamine may be used for the direct precipitation of iron in acid solution, in the presence of aluminium, chromium, cobalt, nickel, and zinc. Copper is precipitated along with the iron, but is easily removed afterwards by treatment with ammonia, in which it is soluble.

The precipitating solution is made by dissolving 6 grams of the amino derivative in 100 grams of water, and may be kept for a week if protected from the light. The solution containing the iron is acidified with concentrated hydrochloric acid and the reagent added until no further precipitation of iron takes place. The precipitate, which is reddish brown in colour, is allowed to settle, washed with twice-normal hydrochloric acid, then with water, ammonia, and water in succession, and finally ignited to Fe2O3.

α-nitroso β-naphthol is a convenient reagent for precipitating iron, particularly if aluminium is present, since this latter metal is not precipitated by the naphthol. The naphthol solution should be made up fresh once a month, as it is rather unstable. For this purpose 4 grams of the solid are dissolved in 150 c.c. of cold glacial acetic acid, and the solution subsequently diluted with an equal quantity of water. The iron should be present in solution as chloride or sulphate, and may be in any state of oxidation - not necessarily all as ferric or all as ferrous metal. The slightly acid solution is mixed with an equal volume of 50 per cent, acetic acid, and an excess of the β-naphthol added. After 6 to 8 hours the solution is filtered, the precipitate, washed first with cold 50 per cent, acetic acid and then with water, dried, and ignited, being weighed as Fe2O3.


Estimation of Iron by Volumetric methods

Volumetric methods are largely employed for the rapid estimation of iron in solution. As a rule, the iron must be present in the ferrous condition, and any ferric iron must first be reduced, the reagents employed varying somewhat with the method of titration. If both the ferrous and ferric contents of a solution are required, the ferrous iron is first determined by titration, then the whole is reduced, and the total iron determined. Subtraction gives the amount of ferric iron originally present.

Bichromate Method

The iron is reduced with stannous chloride or sodium sulphite, and the solution acidified, preferably with sulphuric acid. Potassium bichromate solution is now added from a burette until a spot of the mixture removed on the tip of a glass rod fails to give a blue coloration when mixed with a drop of freshly prepared potassium ferricyanide solution on a white, glazed earthenware tile. The whole of the iron has then been oxidised to the ferric state in accordance with the equation: -

6FeSO4 + K2Cr2O7 + 7H2SO4 = 3Fe2(SO4)3 + K2SO4 + Cr2(SO4)3 + 7H2O.

In the absence of acid, the results obtained are too high, the amount of dichromate required to complete the reaction increasing with dilution of the ferrous salt. With dry ferrous sulphate the result closely approaches the theoretical.

Permanganate Method

The iron is conveniently reduced with zinc or magnesium in the presence of dilute sulphuric acid. A blank experiment should be carried out with the zinc or magnesium alone, in order that a correction may be applied in the event of traces of iron being present as impurity in the metal, and for any carbon which is also liable to affect the titration by reducing a portion of the permanganate. No special indicator is required by this method, for the permanganate is added to the solution, acidified with sulphuric acid until a faint pink colour persists, indicating that the permanganate is now present in slight excess, having oxidised all the iron. Thus: -

10FeSO4 + 2KMnO4 + 8H2SO4 = 5Fe2(SO4)3 + K2SO4 + 2MnSO4 + 8H2O.

The end-point is unstable in the presence of fluorides, but satisfactory results can, in these circumstances, be obtained by addition of fairly concentrated sulphuric or boric acid solutions.

Hydrochloric acid also renders the estimation unreliable unless special precautions are taken, the amount of permanganate used being too great. This has been attributed to the intermediate formation of a higher chloride of manganese during the reduction of the permanganate, which only relatively slowly oxidises the ferrous iron to the ferric condition according to the equation

MnClx+2 + xFeCl2 ⇔ xFeCl3 + MnCl2

until equilibrium is established. Consequently, if the titration is carried out rapidly, more permanganate is added than is theoretically necessary before the ferrous iron is completely, oxidised, and the pink colour indicative of the end-point appears. An excess of the higher manganese chloride remains in solution.

Addition of manganous sulphate, phosphoric acid, and of other substances to the liquid to be titrated has been found to reduce the error.

Stannous chloride

Stannous chloride is often a convenient reagent to use, as it enables a determination to be made of the quantity of ferric iron in the presence of ferrous. The reaction hinges on the fact that stannous chloride reduces ferric chloride to the colourless ferrous salt. Thus: -

2FeCl3 + SnCl2 = 2FeCl2 + SnCl4.

Concentrated hydrochloric acid is added to the iron solution, the whole raised to boiling and quickly titrated with stannous chloride solution until the solution becomes colourless.

To determine the total iron content, the solution is first oxidised with potassium chlorate, and the ferric iron solution thus obtained titrated as above. The difference between two such determinations gives the amount of ferrous iron.

In order that the results should be accurate, especial care must be taken to ensure uniformity of conditions.

Estimation of Iron by Colorimetric Tests

A sensitive test for iron consists in adding a small quantity of hydrazine sulphate to a dilute solution of the iron salt to reduce it to the ferrous condition. Ammonia is now added in excess, and an alcoholic solution of dimethyl glyoxime. The solution is heated to boiling, and then cooled. A faint but detectable rose-red coloration is obtained in the presence of 1 part of iron per 100,000,000 parts of solution.

This method may be applied in the gravimetric estimation of iron. Iron may be detected also by the yellow coloration produced in concentrated hydrochloric acid, 1 part of iron per 100,000 being observable. The colour varies with the concentration of the acid, the maximum intensity occurring with 28 per cent, acid.

Iron in natural waters is frequently estimated colorimetrically by means of the red colour produced by ferric salts with potassium thiocyanate. The iron is best oxidised with nitric acid, as the results appear to be more trustworthy than when oxidation is effected with hydrochloric acid and potassium chlorate. One part of iron in 1,600,000 parts of solution may be detected in this manner.

A delicate reaction for ferrous iron consists in adding a solution of sodium phospho-tungstate, acidified with hydrochloric acid, to the solution suspected of containing a ferrous salt. The whole is rendered alkaline with caustic soda, when a blue coloration is produced if a ferrous salt is present. This reaction is more delicate than that with potassium ferricyanide.

Colorimetric methods are often uncertain in the presence of copper salts, but an accurate method has been worked out, whereby 0.00002 gram of iron can be detected in the presence of 0.2 gram of copper. The method hinges on the violet colour produced when salicylic acid dissolved in acetic acid is added to ferric chloride in the presence of excess of sodium acetate. Under these conditions the depth of colour is proportional to the amount of iron present. The blue or green colour of the copper, which might mask the red produced by the iron, is first removed, where necessary, by addition of dilute potassium cyanide solution.

The red colour produced by ferric iron with acetyl acetone is recommended as the basis of a useful method of estimating the metal colorimetrically.

In acid solution ferrous iron gives an intense blue colour with potassium ferricyanide - the so-called Turnbull's blue. The intensity of the colour is so great that one part of iron in 500,000 parts of solution can be detected.
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