The iodine clock reaction is a classical chemical clock demonstration experiment
to display chemical kinetics in action; it was discovered by Hans Heinrich
Landolt in 1886. Two colourless 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, iodine is
produced: H2O2 + 2I− + 2H+ → I2 + 2H2O
In the second, fast reaction, iodine is reconverted to 2 iodide ions by the
thiosulfate: 2S2O32− + I2 → S4O62− + 2I−
After some time the solution always changes color 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, 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.
Iodate variation An alternative protocol uses a solution
of iodate ion to which an acidified solution of sodium bisulfite is added.
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 to form the dark
blue complex with starch. Persulfate variation
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 color. 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 while ammonium persulfate has a higher solubility and is used instead in
the reaction described in examples from Oxford University.
Chlorate variation 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. Triiodide is present in equilibrium with
iodide anion and molecular iodine: I3− → I− + I2
Chlorate ion oxidizes iodide ion to hypoiodous acid and chlorous acid in
slow and rate-determining step: 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
In this clock the induction period is the time it takes for autocatalytic
process to start after which the concentration of free iodine falls
rapidly as observed by UV/VIS spectroscopy.
See also Clock reaction
Old Nassau reaction References
External links Hydrogen peroxide variation
Potassium iodate variation Potassium iodate variation II
Sodium bisulfite variation with a high-speed camera