Iodometry

Iodometry, also 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.

Free iodine is titrated against standard reducing agent such as sodium thiosulfate solution. Usual reagents are sodium thiosulfate as titrant, starch as an indicator (it forms blue complex with iodine molecules - though polyvinyl alcohol has started to be used recently as well), and an iodine compound (iodide or iodate, depending on the desired reaction with the sample).The color change at the end point is usually blue to colorless.

The principal reaction is the reduction of iodine to iodide by thiosulfate:

I₂ + 2 S₂O₃²⁻ → S₄O₆^2- + 2 I^-

A common and illustrative use of iodometry is the measurement of concentration of chlorine in water. Chlorine in pH under 8 oxidizes iodide to iodine. An overabundance of potassium iodide is added to the known amount of sample in acidic environment (pH < 4, the reaction is not complete in more alkaline pH). Starch is added, forming a blue clathrate complex with the liberated iodine. The blue solution is then titrated with thiosulfate until the blue color vanishes.

Two possible sources of error can influence the outcome of the iodometric titration. One is the air oxidation of acid-iodide solution and the other is the volatility of I₂. The first one can be eliminated by adding an excess of sodium carbonate in the reaction vessel. This displaces the oxygen in the vessel by forming carbon dioxide gas, which blows the air away, and being heavier than air, protects the mixture by remaining on top. The other error can be reduced by using an excess of iodide solution which captures liberated iodine to form triiodide ions, I₃⁻.

Iodometry involves indirect titration of free iodine whereas Iodimetry involves direct titration of liberated iodine.

See also

• Starch indicator
• Redox titration
• Redox indicator
• Cerimetry