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Iodine clock reaction

The Iodine clock reaction or Landolt reaction is a classical chemical clock demonstration experiment to display chemical kinetics in action; it was discovered in 1886 [1]. Two clear solutions are mixed and at first there is no visible reaction, but after a short time delay, the liquid suddenly turns to a shade of dark blue. The iodine clock reaction exists in several variations.


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(aq) + 3 I-(aq) + 2 H+ → I3- + 2 H2O.
In the second, fast reaction, triiodide is reconverted to iodide by the thiosulfate.
I3-(aq) + 2 S2O32-(aq) → 3 I-(aq) + S4O62-(aq)

After some time the solution will rapidly change 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 color caused by the triiodide - starch complex appears.

Anything that accelerates the first reaction will shorten the time until the solution changes color. Increasing the pH, or 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 color to appear.

Iodate variation

An alternative protocol uses a solution of iodate ion (for instance potassium iodate) to which an acidified solution (again with sulfuric acid) of sodium bisulfite is added.

In this protocol, iodide ion is generated by the following slow reaction between the iodate and bisulfite:

IO3- (aq) + 3HSO3- (aq) → I- (aq) + 3HSO4-(aq)

This is the rate determining step. The iodate in excess will oxidize the iodide generated above to form iodine:

IO3- (aq) + 5I- (aq) + 6H+ (aq) → 3I2 + 3H2O (l)

However, the iodine is reduced immediately back to iodide by the bisulfite:

I2 (aq) + HSO3- (aq) + H2O (l) → 2I- (aq) + HSO4-(aq) + 2H+ (aq)

When the bisulphite is fully consumed, the iodine will survive (i.e., no reduction by the bisulfite) to form the dark blue complex with starch.

Persulfate/Peroxydisulphate variation

This clock reaction uses sodium, potassium or ammonium persulfate to oxidise 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:
2I-(aq) + S2O82-(aq) → I2 (aq) + 2SO42-(aq)
And is then removed:
I2 (aq) + 2S2O32-(aq) → 2I-(aq) + S4O62-(aq)

Once all the thiosulphate is consumed the Iodine may form a complex with the starch.[2] Potassium persulfate has a low solute potential, according to documents on the Salters website. Ammonium persulfate has a higher solubility and is used instead in reaction document 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.[3]

iodide anions present in equilibrium with triiodide ion and iodine:
I3- → I- + I2
chlorate ion oxidizes iodide ion to hypoiodous acid and Chlorous acid in slow and rate-determining step:
ClO3- + I- + 2H+ → 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


  1. ^ Landolt, H. Ber. Dtsch. Chem. Ges. 1886, 19, 1317-1365.
  2. ^
  3. ^ André P. Oliveira and Roberto B. Faria (2005). "The Chlorate-Iodine Clock Reaction". J. Am. Chem. Soc. 127 (51): 18022 - 18023. doi:10.1021/ja0570537.
This article is licensed under the GNU Free Documentation License. It uses material from the Wikipedia article "Iodine_clock_reaction". A list of authors is available in Wikipedia.
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