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. The iodine clock reaction exists in several variations, which each involve iodine species (iodide ion, free iodine, or iodate ion) and redox reagents in the presence of starch. 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 due to the formation of a triiodide–starch complex. 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 method starts with a solution of hydrogen peroxide and sulfuric acid. To this a solution containing potassium iodide, sodium thiosulfate, and starch is added. There are two reactions occurring simultaneously in the solution.
In the first, slow reaction, iodine is produced:
- H2O2 + 2 I + 2 H → I2 + 2 H2O
In the second, fast reaction, iodine is reconverted to two iodide ions by the thiosulfate:
- 2 S2O2−3 + I2 → S4O2−6 + 2 I
After some time the solution changes color to a very dark blue, almost black.
When the solutions are mixed, the second reaction causes the iodine to be consumed much faster than it is generated, and only a small amount of iodine is present in the dynamic equilibrium. Once the thiosulfate ion has been exhausted, this reaction stops and the blue colour caused by the iodine – 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.
Aside from using sodium thiosulfate as a substrate, cysteine can also be used.
Iodide from potassium iodide is converted to iodine in the first reaction:
2 I + 2 H + H2O2 → I2 + 2 H2O
The iodine produced in the first reaction is reduced back to iodide by the reducing agent, cysteine. At the same time, cysteine is oxidized into cystine.
2 C3H7NO2S + I2 → C6H12N2O4S2 + 2 I + 2 H
Similar to thiosulfate case, when cysteine is exhausted, the blue color appears.
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:
- IO−3 + 3 HSO−3 → I + 3 HSO−4
This first step is the rate determining step. Next, the iodate in excess will oxidize the iodide generated above to form iodine:
- IO−3 + 5 I + 6 H → 3 I2 + 3 H2O
However, the iodine is reduced immediately back to iodide by the bisulfite:
- I2 + HSO−3 + H2O → 2 I + HSO−4 + 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.
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 + S2O2−8 → I2 + 2 SO2−4
And is then removed:
- I2 + 2 S2O2−3 → 2 I + S4O2−6
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.
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:
- I−3 ⇌ I2 + I
Chlorate ion oxidizes iodide ion to hypoiodous acid and chlorous acid in the slow and rate-determining step:
- ClO−3 + I + 2 H → HOI + HClO2
Chlorate consumption is accelerated by reaction of hypoiodous acid to iodous acid and more chlorous acid:
- ClO−3 + HOI + H → HIO2 + HClO2
More autocatalysis when newly generated iodous acid also converts chlorate in the fastest reaction step:
- ClO−3 + HIO2 → IO−3 + HClO2
In this clock the induction period is the time it takes for the autocatalytic process to start after which the concentration of free iodine falls rapidly as observed by UV–visible spectroscopy.
See also
References
- See:
- Landolt, H. (1886). "Ueber die Zeitdauer der Reaction zwischen Jodsäure und schwefliger Säure" [On the duration of the reaction between iodic acid and sulfurous acid]. Berichte der Deutschen Chemischen Gesellschaft (in German). 19 (1): 1317–1365. doi:10.1002/cber.188601901293.
- Landolt, H. (1887). "Ueber die Zeitdauer der Reaction zwischen Jodsäure und schwefliger Säure [Part 2]" [On the duration of the reaction between iodic acid and sulfurous acid]. Berichte der Deutschen Chemischen Gesellschaft (in German). 20 (1): 745–760. doi:10.1002/cber.188702001173.
- Limpanuparb, T.; Ruchawapol, C.; Sathainthammanee, D. (2019). "Clock Reaction Revisited: Catalyzed Redox Substrate-Depletive Reactions". Journal of Chemical Education. 96 (4): 812–818. Bibcode:2019JChEd..96..812L. doi:10.1021/acs.jchemed.8b00547. S2CID 104370691.
{{cite journal}}: CS1 maint: multiple names: authors list (link) - "Experiment 6: THE RATE LAWS OF AN IODINE CLOCK REACTION" (PDF). Archived from the original (PDF) on 2018-05-17. Retrieved 2018-04-30.
- Hugh Cartwright (2006). "Kinetics of the Persulfate-iodide Clock Reaction" (PDF). 2nd/3rd Year Physical Chemistry Practical Course. Oxford University. Retrieved 25 March 2018.
- 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. PMID 16366551.