Iodine clock reaction
The iodine clock reaction (also known as the Harcourt-Esson reaction or the Harcourt-Essen reaction) is a classical chemical clock demonstration experiment to display chemical kinetics in action; it was discovered by Hans Heinrich Landolt in 1886. Two colorless solutions are mixed and at first there is no visible reaction. After a short time delay, the liquid suddenly turns to a shade of dark blue. The iodine clock reaction exists in several variations. In some variations, the solution will repeatedly cycle from colorless to blue and back to colorless, until the reagents are depleted.
Hydrogen peroxide variation
This reaction starts from a solution of hydrogen peroxide with sulfuric acid. To this is added a solution containing potassium iodide, sodium thiosulfate, and starch. There are two reactions occurring in the solution.
In the first, slow reaction, the triiodide ion is produced:
- H2O2 + 3 I− + 2 H+ → I3− + 2 H2O
In the second, fast reaction, triiodide is reconverted to iodide by the thiosulfate:
- I3− + 2 S2O32− → 3 I− + S4O62−
After some time the solution always changes colour to a very dark blue, almost black.
When the solutions are mixed, the second reaction causes the triiodide ion to be consumed much faster than it is generated, and only a small amount of triiodide is present in the dynamic equilibrium. Once the thiosulfate ion has been exhausted, this reaction stops and the blue colour caused by the triiodide – starch complex appears.
Anything that accelerates the first reaction will shorten the time until the solution changes color. Decreasing the pH (increasing H+ concentration), or increasing the concentration of iodide or hydrogen peroxide will shorten the time. Adding more thiosulfate will have the opposite effect; it will take longer for the blue colour to appear.
In this protocol, iodide ion is generated by the following slow reaction between the iodate and bisulfite:
- IO3− + 3 HSO3− → I− + 3 HSO4−
This is the rate determining step. The iodate in excess will oxidize the iodide generated above to form iodine:
- IO3− + 5 I− + 6 H+ → 3 I2 + 3 H2O
However, the iodine is reduced immediately back to iodide by the bisulfite:
- I2 + HSO3− + H2O → 2 I− + HSO4− + 2 H+
When the bisulfite is fully consumed, the iodine will survive (i.e., no reduction by the bisulfite) to form the dark blue complex with starch.
This clock reaction uses sodium, potassium or ammonium persulfate to oxidize iodide ions to iodine. Sodium thiosulfate is used to reduce iodine back to iodide before the iodine can complex with the starch to form the characteristic blue-black colour.
Iodine is generated:
- 2 I− + S2O82− → I2 + 2 SO42−
And is then removed:
- I2 + 2 S2O32− → 2 I− + S4O62−
Once all the thiosulfate is consumed the iodine may form a complex with the starch. Potassium persulfate is less soluble (cfr. Salters website) while ammonium persulfate has a higher solubility and is used instead in the reaction described in examples from Oxford University.
An experimental iodine clock sequence has also been established for a system consisting of iodine potassium-iodide, sodium chlorate and perchloric acid that takes place through the following reactions.
- I3− → I− + I2
- ClO3− + I− + 2 H+ → HOI + HClO2
Chlorate consumption is accelerated by reaction of hypoiodous acid to iodous acid and more chlorous acid:
- ClO3− + HOI + H+ → HIO2 + HClO2
More autocatalysis when newly generated iodous acid also converts chlorate in the fastest reaction step:
- ClO3− + HIO2 → IO3− + HClO2