THE ESTIMATION OF IRON BY STANDARD POTASSIUM BICHROMATE SOLUTION (Penny’s Method).
Apparatus, Reagents.—For the preparation of the standard solution pure K2Cr2O7 is required, also the iron wire as before, E. SnCl2, 2/5 E. HCl and E. K6Fe2C12N12, freshly prepared and containing no ferrocyanide. A white porcelain plate is used for the indicator tests. For analysis the student may take another sample of the ferric sulphate previously used.
Method, Reactions.—If a solution of K2Cr2O7 be added to a solution of a ferrous salt in presence of a strong free acid, oxidation takes place according to the equation,
6FeCl2 + K2Cr2O7 + 14HCl = 3Fe2Cl6 + Cr2Cl6 + 2KCl + 7H2O
The end point of this reaction is determined by an “external” indicator. A drop of the solution is brought in contact on a white porcelain plate with a drop of K6Fe2C12N12 freshly prepared and containing no K4FeC6N6. A rich blue results if the oxidation has not proceeded far. On adding to the solution more K2Cr2O7 and testing a drop after each addition, the blue changes to a turbid greenish blue, to a grey, and finally to a brown. When the greenish blue has just disappeared the reaction is complete.
When estimating the iron in a ferric salt the reduction is generally brought about, not by zinc, but by stannous chloride, any slight excess of this salt being removed by a few drops of mercuric chloride.
Knowing the volume and value of K2Cr2O7 solution used, the quantity of iron present is easily calculated.
Preparation of the Standard Solution N/10 K2Cr2O7.
As this salt yields 3 atoms of oxygen, which are equivalent to 6 atoms of hydrogen, the normal solution will contain 78 + 104.8 + 112/6 = 49.13 gms. per litre, therefore a solution will contain 4.913 gms. per litre.
Weigh out this quantity of the salt; dissolve it in distilled water, and make up to 1 litre at 16° C. One c.c. of this solution should be equal to .0056 gm. of iron.
Checking the Standard.—Weigh out two portions, each of about .1 gm. clean soft iron wire. Place each portion in a 200 c.c. beaker. To each add 50 c.cs. 5E. HCl. Cover each beaker with a clock glass, and dissolve the wire with the aid of heat. When dissolved remove and rinse the covers. To the hot solution add drop by drop E. SnCl2 with constant stirring. Continue till the yellow colour of the solution has completely disappeared, but avoid excess as far as possible. Dilute to 100 c.cs. Cool quickly and add drop by drop 2/5 E. HgCl2. If no precipitate forms, insufficient SnCl2 was added, and, on the other hand, too copious a precipitate will interfere with the titration. As soon as the precipitate, which should be white, has ceased forming, quickly fill the burette with the N/10 K2Cr2O7 and on a porcelain plate place with a glass rod about a dozen drops of uniform size of K6Fe2Cl2N12, solution. Proceed with the titration, constantly stirring, and removing now and then a drop of the solution on the rod, and bringing it in contact with a drop of the indicator on the porcelain plate. As soon as the bluish green has disappeared, read the burette.
Time will be saved by calculating in advance the probable quantity of solution required, and running in rapidly to within a few c.cs. of this quantity.
Assume that .1022 gm. iron wire is taken and that it requires 18.2 c.cs. N/10 K2Cr2O7. Now .1022 gm. of the wire contain .1022 x .996 = .1018 gm. iron.
1 c.c. N/10 K2Cr2O7 = .1018/18.2 = .005594 gm. Fe
Duplicates should agree within .00002 gm. iron.
The Actual Analysis.—Weigh out 3 gm. anhydrous Fe2(SO4)3. Transfer to a 200 c.c. beaker. Treat exactly the same as when dissolving and reducing the iron wire, and titrate with the N/10 K2Cr2O7 as before.
Repeat the estimation on a fresh portion of the ferric sulphate.
Calculate the results as usual, and repeat the estimation if the duplicates do not agree. The results obtained by this method should agree with those obtained by the permanganate.
Note.—As HCl is a more common solvent than H2SO4, this method is preferable to the permanganate. The solution is more stable and is not affected by rubber, so that the cheaper form of burette may be used. It has the slight disadvantage that an external indicator is required, but with a little practice and careful observation the student will soon learn to sharply distinguish the ‘end point.’