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Description of key information

In contact with water the substance will hydrolyse rapidly and form a pre-equilibrium.  This first step is a first-order reaction with a rate constant of 3.0 s-1 at 20 °C and a pH value of 5. In principle, this reaction can be summarized as I2 + H2O = HOI + I- + H+. As hydronium ions are formed in this reaction it can be expected that natural environmental conditions support the formation of hypoiodous acid (HOI) and iodide. In the second step of the hydrolysis HOI disproportionates and forms iodide and iodate via a mechanism not fully understood. This step is considered as a second-order reaction and is clearly slower taking at least tens of minutes depending on the prevalent conditions. Additionally, the findings in the different studies indicate that the initial reaction of the second disproportionation is relative fast, but then after about one hour the gradient of the rate constants flatten reaching for steady-state conditions in solutions at pHs between 6 and 8. Anyhow, it can be assumed that at that time the major partition of molecular iodine is already degraded. The overall hydrolysis can be summarized as 3 I2 + 3 H2O = IO3- + 5 I- + 6 H+.

Since the first disproportionation and the fast initial step of the second equilibrium reaction play a decisive role in the environmental fate of iodine and adverse effects of HOI in the environment cannot be excluded the rate constant of the first reaction is considered for the risk assessment.  

Key value for chemical safety assessment

Half-life for hydrolysis:
0.01 min
at the temperature of:
12 °C

Additional information

The hydrolysis of iodine has been in the focus of the scientific community for several decades. From various studies it has become clear that it is a very complex system consisting of competing partial reactions forming two equilibria in principle and a vast number of different iodine species which are formed and consumed in different pathways.The overall hydrolysis can be summarized as 3 I2 + 3 H2O = IO3- + 5 I- + 6 H+. Hypoiodous Acid (HOI) seems to be a key intermediate in the whole system.

In 1955 Allen et al. examined the hydrolysis of iodine and determined the equilibria constant for the formation of hypoiodous acid for two different temperatures (K= 5.4 * 10-13, 25 °C and K= 0.49x10-13, 1.6 °C) by photometric determination of the iodine and triiodide concentration in the solution. Based on previous observations of Liebhafsky (Liebhafsky, 1931) and Morgan (Morgan, 1954) they postulated a reaction system for the formation of hypoiodous acid consisting of different competing partial reactions.

I2 + H2O = H2OI+ + I-

I2 + H2O = HOI + H+ + I-

I2 + I- = I3-

Eigen and Kustin in 1962 studied this reaction system by a temperature-jump relaxation technique. They extended the system of competing partial reactions for the formation of hypoiodous acid with additional potential routes and iodine species and were able to propose a general mechanism. Additionally, they were able to determine a rate constant for the first pre-equilibrium (formation of HOI) at a pH value of 5 (k= 3.0 s-1).

In 1980 Thomas and co-workers studied the hydrolysis of iodine and disproportionation of hypoiodous acid to iodide and iodate by a combination of photometrical and potentiometrical analytical techniques for pH values between 7 and 12. Their experimentally found rate constants were in good agreement to their numerical integration of the theoretical reaction kinetics. They determined the extent of the disproportionation by comparing the initial iodine concentration and the concentration of the formed iodine species. About 25 % of the injected iodine remained after 4 days at pH 7, 1% remained after 14 hours at pH 8 and about 0.1% after 2 hours at pH 9. For higher pH values the ratio of initial iodine and formed iodide concentrations approached the theoretical limit, but did not reached it.

Truesdale et al. examined the disproportionation of hypoiodous acid at pHs between 7 and 13 without addition of iodide. Based on their results they were able to develop a model of a second order kinetics depending on the hypoiodous acid concentration influenced by two competing pre-equilibria involving iodine and hypoiodous acid as well as hypoiodous acid and hypoiodite. Because the rate maximum appeared at ca. pH 9.5 he assumed a central role of hypoiodous acid in the disproportionation mechanism as this species pre-dominates the system at this pH value.

Such as Thomas et al. Truesdale observed the formation of a steady-state condition of the disproportionation for pHs between 6 and 8 after a reasonably fast initial reaction. The extent of the initial reaction decreases with decreasing pH. While at pH 6.82 only 21% of the iodine is converted the extent increases to 93% at pH 8.33 within the first hour (Truesdale, 1995). Afterwards a phase of a distinctly decreased reaction rate is established which is indicated by a brown color of the solution that can remain for days or even weeks at 25° C. Likewise, Luther and Sammet (Luther, 1905) reported a period of 49 days to reach equilibrium at pH 6.2. Sugawara and Terada (Sugawara, 1958) have observed that iodate was still being produced in sea water after 25 days.

In general the iodine - hypoiodous acid equilibrium is very fast as Eigen and Kustin showed, and takes a fraction of a second at a temperature of 20 °C and pH value of 5 to adjust. Thus, it can be expected that most of the iodine in the environment is rapidly transformed as the natural pH value actually promotes the pre-equilibrium reaction. In contrast, the step of the disproportionation of hypoiodous acid to iodide and iodate is much slower, taking at least tens of minutes to reach completion at 25 °C. Additionally, a strong connection with the pH value for this reaction is observable reaching higher reaction rates for readily alkaline environment. However, about 90% of all iodine found in natural freshwater occurs as iodide (FOREGS Geochemical Baseline Programme (FGBP) database).


Liebhafsky HA (1931). Reactions involving hydrogen peroxide, iodine and iodate ion. IV. The Oxidation of iodine to iodate ion by Hydrogen Peroxide, J. Am. Chem. Soc., 53, 2074-2090.

Luther R (1905). Sammet GV, Z. Elektrochem. 11, 293.

Morgan KJ (1954). Some reactions of inorganic iodine compounds., Quart. Rev., 8, 123-146.

Truesdale VW, Canosa-Mas C (1995). Kinetics of Disproportionation of Hypoiodous Acid in Phosphate and Borate Buffer at pH<8.5 modelled using Iodide Feedback, J.Chem. Soc. Faraday Trans., 91(15), 2269 -2273.

Sugawara K, Terada K (1958). Oxidised iodine in sea water, Nature, 182, 250.