page_title: Iodometry
section_title: Basic principles
hierarchical_section_title: Iodometry / Basic principles
caption_attribution_description: English: Standard solution of iodine sealed in an ampoule for iodometric analysis
context_page_description: Iodometry, known as iodometric titration, is a method of volumetric chemical analysis, a redox titration where the appearance or disappearance of elementary iodine indicates the end point.
  Note that iodometry involves indirect titration of iodine liberated by reaction with the analyte, whereas iodimetry involves direct titration using iodine as the titrant.
  Redox titration using sodium thiosulphate, Na₂S₂O₃ as a reducing agent is known as iodometric titration since it is used specifically to titrate iodine. The iodometric titration is a general method to determine the concentration of an oxidising agent in solution. In an iodometric titration, a starch solution is used as an indicator since it can absorb the I₂ that is released. This absorption will cause the solution to change its colour from deep blue to light yellow when titrated with standardised thiosulfate solution. This indicates the end point of the titration. Iodometry is commonly used to analyze the concentration of oxidizing agents in water samples, such as oxygen saturation in ecological studies or active chlorine in swimming pool water analysis.
context_section_description: To a known volume of sample, an excess but known amount of iodide is added, which the oxidizing agent then oxidizes to iodine. Iodine dissolves in the iodide-containing solution to give triiodide ions, which have a dark brown color.  The triiodide ion solution is then titrated against standard thiosulfate solution to give iodide again using starch indicator:
  I₃⁻ + 2 e⁻ ⇌ 3 I⁻ (Eᵒ = + 0.5355 V)
  Together with reduction potential of thiosulfate:
  S₄O₆²⁻ + 2 e⁻ ⇌ 2 S₂O₃²⁻ (Eᵒ = + 0.08 V)
  The overall reaction is thus:
  I₃⁻ + 2 S₂O₃²⁻ → S₄O₆²⁻ + 3 I⁻ (Ereaction = + 0.4555 V)
  For simplicity, the equations will usually be written in terms of aqueous molecular iodine rather than the triiodide ion, as the iodide ion did not participate in the reaction in terms of mole ratio analysis. The disappearance of the deep blue color is, due to the decomposition of the iodine-starch clathrate, marks the end point.
  The reducing agent used does not necessarily need to be thiosulfate; stannous chloride, sulfites, sulfides, arsenic(III), and antimony(III) salts are commonly used alternatives. at higher pH (> 8)).
  At low pH, the following reaction might occur with thiosulfate:
  S₂O₃²⁻ + 2 H⁺ → SO₂ + S + H₂O
  Some reactions involving certain reductants are reversible at certain pH, thus the pH of the sample solution should be carefully adjusted before the performing the analysis. For example, the reaction:
  H₃AsO₃ + I₂ + H₂O → H₃AsO₄ + 2 H⁺ + 2 I⁻
  is reversible at pH < 4.
  The volatility of iodine is also a source of error for the titration, this can be effectively prevented by ensuring an excess iodide is present and cooling the titration mixture. Strong light, nitrite and copper ions catalyzes the conversion of iodide to iodine, so these should be removed prior to the addition of iodide to the sample.
  For prolonged titrations, it is advised to add dry ice to the titration mixture to displace air from the erlenmeyer flask so as to prevent the aerial oxidation of iodide to iodine.  Standard iodine solution is prepared from potassium iodate and potassium iodide, which are both primary standards):
  IO₃⁻ + 8 I⁻ + 6 H⁺ → 3 I₃⁻ + 3 H₂O
  Iodine in organic solvents, such as diethyl ether and carbon tetrachloride, may be titrated against sodium thiosulfate dissolved in acetone.
caption: Iodine standard solution, sealed in an ampoule for iodometric analysis
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