Iodometry

The Iodometrie is a method of chemical analysis that is used for the quantitative determination of various substances. It belongs to the titrimetric analysis method and is based on the conversion of iodide ions (I ^- ) into iodine (I ₂ ) or vice versa, i.e. on the following equilibrium reaction :

${\displaystyle \mathrm {2I^{-}\rightleftharpoons I_{2}+2e^{-}} }$

Iodometry can be used to quantitatively determine analytes that have an I ^- oxidizing effect as well as I ₂ reducing analytes. Iodometry must be clearly differentiated from iodine atomometry, the latter using the redox properties of iodate ions.

Determination of reducing analytes and preparation of the standard iodine solution

Reducing effect analytes can I ₂ reduce to iodide ions. The higher the amount of analyte in the sample, the more iodine can be reduced. Such analytes can be determined either by direct titration or by back titration.

Direct titration

With direct titration, the sample is titrated directly with iodine standard solution. From the volume consumption of standard iodine solution up to the equivalence point, the amount of analyte can be calculated back using the stoichiometric laws. Direct titration is recommended if the analyte reacts completely and quickly with iodine. This enables an exact recognition of the equivalence point.

Determination by means of back titration

The determination of reducing analytes can also be carried out using a back titration . For this purpose, the sample is mixed with a defined excess (i.e. known concentration and known volume) of standard iodine solution. After leaving it to stand for a short time, the unreacted residual amount of iodine can be determined by titration with sodium thiosulphate solution. Such a back titration is advisable if the analyte only reacts slowly with iodine and therefore cannot be titrated with it directly.
An example of this is the determination of mercury (I) chloride , which reacts with an iodine / iodide solution to form a complex compound :

${\displaystyle \mathrm {Hg_{2}Cl_{2}+I_{2}+6\ I^{-}\longrightarrow 2\ [HgI_{4}]^{2-}\ +2\ Cl^{-}} }$

Mercury (II) salts can be determined in a similar way .

Preparation of the standard solution and determination of the titer

Iodine itself is only sparingly soluble in water. Therefore, when preparing the standard solution, the iodine is added to a potassium iodide solution. In this salt solution, iodine dissolves much better with the formation of triiodide ions:

${\displaystyle \mathrm {I_{2}+I^{-}\rightleftharpoons I_{3}^{-}} }$

Since the exact weighing of iodine is difficult due to the high vapor pressure or the volatile character on the weighing pan, the iodine standard solution is often also prepared using potassium iodate and potassium iodide. For this purpose, the exact required mass of the easily weighable potassium iodate is added to a solution with an excess of potassium iodide. After acidification, the desired iodine content is formed in a comproportionation reaction . The standard iodine solution is unstable, and the degradation of iodine is particularly caused by exposure to light. In order to better protect against light radiation, the solution is best kept in amber glass containers. The actual iodine content in the standard solution must be checked at regular intervals. For this purpose, a precisely weighed mass of arsenic (III) oxide solution can be used as the base substance , which is quantitatively converted with the iodine to arsenates or arsenic acid. Because of the toxicity of arsenic compounds, the actual content can alternatively be determined with sodium thiosulfate pentahydrate as a solid or with a sodium thiosulfate solution of known content.

Determination of oxidizing analytes

Analytes with an oxidizing effect oxidize iodide ions to iodine, the resulting amount of iodine being a measure of the amount of analyte. An excess of iodide ions is added to the sample in order to rule out a corresponding deficiency that would limit the formation of iodine. Under these conditions, the analyte can react completely with iodide ions; a residual amount of the latter remains in the sample mixture. The higher the amount of analyte, the higher the amount of iodine produced. The exact ratio of the amount of analyte to the amount of iodine produced varies depending on the analyte and can be read from the ratio of coefficients in the reaction equation. Example for the iodometric determination of Cu ²⁺ :

${\displaystyle \mathrm {2Cu^{2+}+4\ I^{-}\longrightarrow 2\ CuI+\ I_{2}} }$

The amount of substance of I ₂ is exactly half the amount of substance of the analyte (Cu ²⁺ ). The amount of iodine produced in the chemical reaction is determined quantitatively by titration. For this purpose, the sample, which now contains iodine, is titrated with sodium thiosulphate standard solution (Na ₂ S ₂ O ₃ ) in the weakly acidic pH range (in the alkaline range the thiosulphate is oxidized to sulphate ). The iodine is converted back into iodide ions:

${\displaystyle \mathrm {I_{2}+2\ S_{2}O_{3}^{2-}\longrightarrow 2\ I^{-}+\ S_{4}O_{6}^{2-}} }$

At the equivalence point, this reaction has just ended, and from the volume consumption of standard solution up to this point, conclusions can be drawn about the amount of iodine that was present. The higher the consumption of standard solution, the higher the amount of iodine and ultimately the higher the amount of substance of the analyte. Since the analyte is not titrated directly, iodometry for the determination of oxidizing analytes is an indirect titration.

Indication of the equivalence point

Depending on which type of analyte is determined, iodine is produced at the equivalence point or the last remaining iodine is being chemically converted. Although dissolved iodine is slightly yellow in color, the intrinsic color is too weak for an exact endpoint detection (formation / disappearance of the yellow color). Instead, a few drops of starch solution are added to the sample shortly before the equivalence point is reached. The intense blue color of the iodine-starch complex now allows an exact recognition of the equivalence point.

Importance of iodometry

Iodometry is a universal method, as many analytes have reducing or oxidizing properties. However, this is also where the method's greatest problem lies. The determination is often interfered with by other substances contained in the sample, which can also have reducing or oxidizing properties on iodine or iodide ions. An example of an important field of application is the determination of the iodine number . This can be used to measure the number of double bonds in a long-chain alkene or an unsaturated fatty acid .

See also

Manganometry

Individual evidence

1. ^ Brockhaus ABC Chemie , VEB FA Brockhaus Verlag Leipzig 1965, pp. 605-606.
2. ↑ Jander / Jahr / Knoll: Maßanalyse, Göschen Collection Volume 221, Walter de Gruyter Berlin 1966, p. 117.
3. ↑ Otto-Albrecht Neumüller (Ed.): Römpps Chemie-Lexikon. Volume 3: H-L. 8th revised and expanded edition. Franckh'sche Verlagshandlung, Stuttgart 1983, ISBN 3-440-04513-7 , pp. 1916-1917.

literature

• G. Schwedt: Analytical chemistry . 2nd Edition. Wiley-VCH, 2008, ISBN 978-3-527-31206-1 .