Iodine

Hydrolysis

Link to relevant study record(s)

Reference

Reference 1

Endpoint:

hydrolysis

Type of information:

experimental study

Adequacy of study:

weight of evidence

Reliability:

2 (reliable with restrictions)

Rationale for reliability incl. deficiencies:

other: Study well documented, meets generally accepted scientific principles, acceptable for assessment in context of a weight of evidence approach

Qualifier:

no guideline followed

Principles of method if other than guideline:

The concentrations of iodine and triiodide were analysed spectrometrically for ten different buffer solutions (pH between 2.04 and 5.67) and two different temperatures (1.6 °C and 25 °C).

GLP compliance:

no

Details on sampling:

In the flasks with the pre-treated iodine solutions the aqueous phase were replaced with a buffer solution. For the more acidic solutions (pH 2.04 and 3.04) the air in the flask was removed by bubbling nitrogen through the solutions in order to minimize oxidation of solved iodide. All flasks were covered with aluminium foil to exclude light, shaken vigorously for several minutes and afterwards rotated for one hour in a thermostat. Then the aqueos phase was removed and placed in a one-cm.ground-stoppered quartz cell.
The spectrometer was equiped with a water jacket to control the temperature of the quartz cell. In three experiments after performing the measurements at 25 °C the temperature was reduced to 1.6 °C to determine the influence of the temperature. To prevent condensation on the surface of the cell it was flushed by a dry, chilled nitrogen flow.

Buffers:

Buffer solutions were prepared from perchloric acid, sodium perchlorate, sodium acetate and water. The ion strength in all solutions was µ=0.01 M and the pH value ranged from 2.04 to 5.67.

Details on test conditions:

The used water was redestilled from an alkaline premanganate solution. Carbon tetrachloride was of Spectro Grade and redestilled before use. All other chemicals were of a c.p. or reagent grade.

Transformation products:

not measured

Details on hydrolysis and appearance of transformation product(s):

The hydrolysis of iodine is a very complex system of competing reactions with two different equilibria in which various iodine species are formed and transformed in different pathways. The overall reaction can be summarized as 3 I2 + 3 H2O = IO3- + 5 I- + 6 H+. Hypoiodous acid (HOI) seems to be the key intermediate in the whole system.

Remarks on result:

not determinable

Details on results:

Based on the determined extinction coefficients and the observed optical densities the concentrations of iodine and triiodide in the buffer solutions were calculated.

In the fast equilibrium of the hydrolysis of iodine in water different competing reactions occur.

I₂(aq) + H₂O = H₂OI⁺ + I^-

I₂(aq) + H₂O = HOI + H⁺ + I^-

I₂ (aq) + I^- = I₃^-

By combination of the equilibrium equitations a term can be set up which expresses the dependency of the equilibrium constants from the pH value and the correlation with the concentration of the triiodide. By plotting (K₁ + K₂ /[H⁺] ) against the reciproc concentration of hydronium ions a linear correlation can be found. The slope of the line is equal to K₂ and the interception is equal to K₁.

In the series of experiments it could be found that K₁ is much smaller than K₂ and could therefore be neglected in the calculation.

For K₂ a value of (5.40 +/- 0.25) E-13 for 25 °C and (0.49 +/- 0.03)E-13 for 1.6 °C could be determined. An exact determination of K₁ was not possible, only an upper limit of 1E-10 could be established.

Validity criteria fulfilled:

not applicable

Remarks:

When study was conducted no guideline was established.

Conclusions:

This study is not conducted according to a standard test guideline, but meets the generally accepted scientific principles. The findings of this study are used in several further studies. Therefore they can be considered as accepted and peer-reviewed by the scientific community.

Each study for itself does not fulfill the requirements of this endpoint, therefore a combined weight of evidence approach is applied. Further the application of the standard test guideline OECD 111 is considered to be not adequate due to the complex system of hydrolysis of iodine. The standard protocol tends to the identification of the hydrolysis products and the rate of hydrolysis as a function of the pH. As the hydrolysis of iodine is a two-step equilibrium reaction with numerous competing reactions and iodine species an adaptation of the test guideline would be appropriate. Based on available and evaluated literature data we tried to draw a complete picture of the behavior of molecular iodine in water.

In this case the first ultra-fast equilibrium reaction plays a decisive role for the concentration of molecular iodine in water which makes the use of special analytical techniques necessary. Eigen and Kustin analyzed the reaction in acidic media and from the stoichiometry it can be concluded that at environmental pH the reaction will be promoted due to the removal of the formed H+-ions. Thus, a fast and nearly complete transformation can be assumed. The findings of Thomas et al. and Truesdale et al. indicate that the second reaction is of second-order kinetics at pHs between 6 and 10 and therefore depending on the concentration of iodide and hypoiodous acid. Additionally, it is obvious from the observations that, in principle, the initial phase of the second-step disproportionation at environmental pH is fast. Then the gradient of the reaction rate flattens reaching for a steady-state at pHs between 6 and 8 which could remain for several weeks. At the end of the initial phase which lasts for about one hour depending on environmental pH the major partition of iodine is disproportionated to iodide and iodate.
Thus, the presented weight of evidence is considered to be adequate for assessing the environmental risk and describing the fate and behavior of molecular iodine in the aquatic compartment at different pH.

Executive summary:

Allen and Keefer were able to determine the equilibrium constant for the fast hydrolysis fo iodine by spectrophotometric analysis. A value of 5.40 E-13 (25 °C) and 0.49E-13 (1.6 °C) could be established. Furthermore they were able to establish an upper limit for the equilibrium constant of 1E-10 for the reaction I₂(aq) + H₂O = H₂OI⁺ + I^- for 25 °C.

Reference 2

Endpoint:

hydrolysis

Type of information:

experimental study

Adequacy of study:

weight of evidence

Reliability:

2 (reliable with restrictions)

Rationale for reliability incl. deficiencies:

other: see 'Remark'

Remarks:

Study documentation acceptable, meets generally accepted scientific principles and author can be considered as expert for this topic of examinations of ultra-fast reaction kinetics, acceptable for assessment in context of a weight of evidence approach

Qualifier:

no guideline followed

Principles of method if other than guideline:

By using the temperature jump relaxation method, the kinetics of the three different halogens chlorine, bromine and iodine were studied. The temperature rise (about 10 °C) was caused by a high voltage pulse through the system and perturb the equilibrium within 5E-7 sec. To avoid decomposition of iodine in solution solid gold electrodes have been used.

Basic principle of the relaxation technique:
A reaction system is in equilibrium. By a rapid pulse (rapid increase of temperature) this equilibrium is perturbed and the system tries to
adapt to the new conditions. The time for this adaptation (relaxation time) can be measured. In combination with the equilibrium constant for the newly formed equilibrium the rate coefficients for the reaction can be calculated.

GLP compliance:

no

Details on sampling:

Different samples were prepared from weighed out portions of iodine. The concentrations were 0.00005 mol/L, 0.001 mol/L and 1.0E-8 mol/L. The concentrations were controlled by the comparison of the measured and calculated concentrations of the triiodide ion. A good agreement was found. Additionally, by this approach the stability of the solutions could be controlled. Assuming no further reactions occur in the solution and any decomposition of iodine produces iodide, the analysis of triiodide provides a very sensitive test for the hydrolysis.
Because of the limited solubility of iodine in water the range of the concentrations was restricted.

Buffers:

no buffers were used

Details on test conditions:

The experiments were performed in pure water at an ionic strengths of 0.1 M and an equilibrium temperature of 20 °C. Thermodynamic equilibrium constants were corrected by appropriated activity coefficients, where necessary. Furthermore the accuracy of the experiments was limited by a slight electrolytic formation of iodide which altered the concentration of the triiodide during the experiments. The relative error of the rate coefficients is about 10%.

Transformation products:

not measured

Details on hydrolysis and appearance of transformation product(s):

The hydrolysis of iodine is a very complex system of competing reactions with two different equilibria in which various iodine species are formed and transformed in different pathways. The overall reaction can be summarized as 3 I2 + 3 H2O = IO3- + 5 I- + 6 H+. Hypoiodous acid (HOI) seems to be the key intermediate in the whole system. As the stable transformation products are already known from former findings, no analytical identification was performed in this study.

pH:

5

Temp.:

20 °C

Hydrolysis rate constant:

ca. 3 s-1

DT50:

ca. 0.005 min

Type:

(pseudo-)first order (= half-life)

The rate constant for the hydrolysis reaction is calculated from the equilibrium constant (K_hyd= k_forw/k_backw = 4.3 E-13 (Allen TL and Keefer RM (1955). J Am Chem Soc, 77, 2957 -2960, recalculated for experimental conditions) and the rate constant of the backward reaction. Therefor the reciproc values of the measured relaxation times are plotted against ([IOH][H+]+[IOH][I-]+[I-][H+]) and the slope is equal to the rate constant for the backward reaction.

Validity criteria fulfilled:

not applicable

Remarks:

study presented in context of a weight of evidence

Conclusions:

When this study was conducted no guideline on hydrolysis was established. Since released it is one of the most cited reference in this topic and its findings are used in many cases for further investgations. Thus, it can be considered as accepted and peer-reviewed by the scientific community. Furthermore the author can be considered as expert because of receiving the Nobel prize in 1967 for the examination of ultra-fast reaction kinetics by relaxation techniques.

Each study for itself does not fulfil the requirements of this endpoint, therefore a combined weight of evidence approach is applied. Further the application of the standard test guideline OECD 111 is considered to be not adequate due to the complex system of hydrolysis of iodine. The standard protocol tends to the identification of the hydrolysis products and the rate of hydrolysis as a function of the pH. As the hydrolysis of iodine is a two-step equilibrium reaction with numerous competing reactions and iodine species an adaptation of the test guideline would be appropriate. Based on available and evaluated literature data we tried to draw a complete picture of the behavior of molecular iodine in water.

In this case the first ultra-fast equilibrium reaction plays a decisive role for the concentration of molecular iodine in water which makes the use of special analytical techniques necessary. Eigen and Kustin analysed the reation in acidic media and from the stoichiometry it can be concluded that at environmental pH the reaction will be promoted due to the removal of the formed H+-ions. Thus, a fast and nearly complete transformation can be assumed. The findings of Thomas et al. and Truesdale et al. indicate that the second reaction is of second-order kinetics at pHs between 6 and 10 and therefore depending on the concentration of iodide and hypoiodous acid. Additionally, it is obvious from the observations that, in principle, the inital phase of the second-step disproportionation at environmental pH is fast. Then the gradient of the reaction rate flattens reaching for a steady-state at pHs between 6 and 8 which could remain for several weeks. At the end of the inital phase which lasts for about one hour depending on environmental pH the major partition of iodine is disproportionated to iodide and iodate.
Thus, the presented weight of evidence is considered to be adequate for assessing the environmental risk and describing the fate and behavior of molecular iodine in the aquatic compartment at different pH.

Executive summary:

In this study of Eigen and Kustin the fast equilibrium of the hydrolysis reaction of iodine has been examined by the temperature jump relaxation method. This method is suitable to analyse even the fastest equilibrium reactions known (e.g H3O+(aq) + OH-(aq) -> H2O(l)). The overall stoichiometry of the reaction for a pH value between 2 and 7 is I2 + H2O = HOI + I- + H+. Depending on the initial concentration of iodine a significant portion of the iodide produced reacts with iodine and forms triiodide. Because the high molar absorptivity of triiodide is a good indicator for changes in the equilibrium and therefore used to photospectrometrically analyse changes in the equilibrium.

 

For the experiments defined amounts of iodine were solved in pure water without a buffering system and the established equillibrium was disturbed by a high voltage pulse which increased the temperature within 5E-7 s in the solution by about 10° C. The shift of the equilibrium was followed photometrically by measuring the changes in the concentration of triiodide ion at a wavelength of 350 nm. Additionally, the analysis of triiodide ions provides a very sensitive test for the stability of the solution before the disturbance. Since hydronium-ions arise stoichiometrically in the hydrolysis, the compositions were also checked by pH measurements.

The relation between the relaxation time and the rate constants is described by 1/t = kforw+kbackw([IOH]eq[H+]eq+[IOH]eq[I-]eq+[I-]eq[H+]eq). By plotting 1/t against ([IOH]eq[H+]eq+[IOH]eq[I-]eq+[I-]eq[H+]eq, the effect of the concentration product and pH on the relaxation time is shown. The ordinate intercept is equal to kforw and the slope is equal to kbackw. Because kforw is much smaller than kbackw ([IOH]eq[H+]eq+[IOH]eq[I-]eq+[I-]eq[H+]eq) Eigen und Kustin obtained the value of kforw by deriving from the equilibrium equation Khyd = kforw/kbackw.

Reference 3

Endpoint:

hydrolysis

Type of information:

experimental study

Adequacy of study:

weight of evidence

Reliability:

2 (reliable with restrictions)

Rationale for reliability incl. deficiencies:

other: Study well documented, meets generally accepted scientific principles, acceptable for assessment in context of a weight of evidence approach

Qualifier:

no guideline followed

Principles of method if other than guideline:

Spectrophotometric measurement of concentrations of iodine and triiodide in dilute aqueous solutions of iodine. Additionally, the iodide concentration is determined by an iodide ion sensitive electrode. From these concentrations and the total amount of iodine oxidant in the system, the concentration of HOI is calculated.

GLP compliance:

no

Details on sampling:

Defined amounts of iodine were dissolved in 0.5 mL of diethylether which was injected into 100 mL of Nitrogen-saturated and temperated (25 °C) buffer solution. The solutions were virgorously mixed and positioned in the spectrometer within 60 seconds after injection.

Buffers:

All buffer solutions were of an ionoíc strength of 0.15 mol/L. All chemicals used were of reagent grade quality.
for pH 6, 7 and 8: combination of CO2-free NaOH and NaH2PO4
for pH 9: boric acid
for pH 10: NaHCO3

Transformation products:

not measured

Details on hydrolysis and appearance of transformation product(s):

The hydrolysis of iodine is a very complex system of competing reactions with two different equilibria in which various iodine species are formed and transformed in different pathways. The overall reaction can be summarized as 3 I2 + 3 H2O = IO3- + 5 I- + 6 H+. Hypoiodous acid (HOI) seems to be the key intermediate in the whole system.

pH:

7

Temp.:

25 °C

Hydrolysis rate constant:

ca. 0.001 min-1

DT50:

ca. 1 243 min

Type:

second order

Remarks on result:

other: calculated for PNECfreshwater

pH:

8

Temp.:

25 °C

Hydrolysis rate constant:

ca. 0.003 min-1

DT50:

ca. 320 min

Type:

second order

Remarks on result:

other: calculated for PNECfreshwater

pH:

9

Temp.:

25 °C

Hydrolysis rate constant:

ca. 0.003 min-1

DT50:

ca. 339 min

Type:

second order

Remarks on result:

other: calculated for PNECfreshwater

pH:

10

Temp.:

25 °C

Hydrolysis rate constant:

ca. 0.001 min-1

DT50:

ca. 932 min

Type:

second order

Remarks on result:

other: calculated for PNECfreshwater

The results for the rate constants at the different pH values were determined by plotting reaction time against the sum of the concentrations of iodine, triiodide, hypoiodous acid and hypoiodite (slope = rate constant) and by deriving from the numerical integration of the kinetic equitations for the hydrolysis system. The results for the pH between 7 and 10 were consistent. For pH 6 only insufficient data were available. For pH 10 only data of the ion selective electrode were used. All rate constants in Table 1 are shown in L mol^-1 sec^-1.

Table 1: rate constants for hydrolysis at 25.0 °C and 0.15 mol L^-1ionic strength

pH	From slope of concentration - time plot	From numerical integration
7	0.9 E+2	1.0 E+2
8	3.5 E+2	3.0 E+2
9	3.3 E+2	4.5 E+2
10	1.2 E+2	1.4 E+2

Validity criteria fulfilled:

not applicable

Remarks:

data presented in context of a weight of evidence

Conclusions:

Although the presented study is not conducted according a standard test guideline, it meets generally accepted scientific principles and it is well documented.

Each study for itself does not fulfill the requirements of this endpoint, therefore a combined weight of evidence approach is applied. Further the application of the standard test guideline OECD 111 is considered to be not adequate due to the complex system of hydrolysis of iodine. The standard protocol tends to the identification of the hydrolysis products and the rate of hydrolysis as a function of the pH. As the hydrolysis of iodine is a two-step equilibrium reaction with numerous competing reactions and iodine species an adaptation of the test guideline would be appropriate. Based on available and evaluated literature data we tried to draw a complete picture of the behavior of molecular iodine in water.

In this case the first ultra-fast equilibrium reaction plays a decisive role for the concentration of molecular iodine in water which makes the use of special analytical techniques necessary. Eigen and Kustin analyzed the reaction in acidic media and from the stoichiometry it can be concluded that at environmental pH the reaction will be promoted due to the removal of the formed H+-ions. Thus, a fast and nearly complete transformation can be assumed. The findings of Thomas et al. and Truesdale et al. indicate that the second reaction is of second-order kinetics at pHs between 6 and 10 and therefore depending on the concentration of iodide and hypoiodous acid. Additionally, it is obvious from the observations that, in principle, the initial phase of the second-step disproportionation at environmental pH is fast. Then the gradient of the reaction rate flattens reaching for a steady-state at pHs between 6 and 8 which could remain for several weeks. At the end of the initial phase which lasts for about one hour depending on environmental pH the major partition of iodine is disproportionated to iodide and iodate.
Thus, the presented weight of evidence is considered to be adequate for assessing the environmental risk and describing the fate and behavior of molecular iodine in the aquatic compartment at different pH.

Executive summary:

Thomas et. al. studied the disproportionation of dilute aqueous solutions between pH 7 and 10 by photometric and potentiometric analytical methods. The determined experimental rate constants are consistent with rate constants which were gained by numerical intergration of the reaction kinetics.

Reference 4

Endpoint:

hydrolysis

Type of information:

read-across from supporting substance (structural analogue or surrogate)

Adequacy of study:

supporting study

Reliability:

2 (reliable with restrictions)

Rationale for reliability incl. deficiencies:

study well documented, meets generally accepted scientific principles, acceptable for assessment

Qualifier:

no guideline followed

Principles of method if other than guideline:

All investigations were run at a temperature of 25.0℃. At the appropriate time sub-samples (0.400 mL) of each reaction mixture were placed in 40.0 mL of 0.15 mol/L sodium dihydrogen phosphate buffer at pH 6.7. The absorbances of these solutions at 285 nm were measured with a Pye-Unicam SP-500 series I1 spectrophotometer equipped with 10 cm cuvettes.

GLP compliance:

not specified

In four phosphate solutions at pHs between 6.82 and 8.33, a rapid initial reaction led to the establishment of the sluggish phase within ca. 1 h. During the first hour the reaction proceeded to only 21% of its full extent at pH 6.82, whereas at pH 8.33 it attained 93% of its full extent.

Validity criteria fulfilled:

not applicable

Conclusions:

While at pH 6.82 only 21% of the iodide [∑I2] is converted the extent increases to 93% at pH 8.33 within the first hour

Executive summary:

The kinetics of disproportionation of hypoiodous acid has been studied in phosphate and borate buffers at pH <8.5, and without added iodide. In this region, after an initial rapid phase of reaction lasting tens of minutes, the reaction becomes very sluggish sometimes continuing for weeks at 25°C.

In four phosphate solutions at pHs between 6.82 and 8.33, a rapid initial reaction led to the establishment of the sluggish phase within ca. 1 h. During the first hour the reaction proceeded to only 21% of its full extent at pH 6.82, whereas at pH 8.33 it attained 93% of its full extent.

Reference 5

Endpoint:

hydrolysis

Type of information:

experimental study

Adequacy of study:

supporting study

Reliability:

2 (reliable with restrictions)

Rationale for reliability incl. deficiencies:

study well documented, meets generally accepted scientific principles, acceptable for assessment

Qualifier:

no guideline followed

Principles of method if other than guideline:

The approach detailed by Truesdale and Moore (1992) was used throughout.
Truesdale V. W. and Moore, R. M. (1992) Further studies on the chemical reduction of molecular iodine added to seawater. Mar. Chem. 40, 199-213.

GLP compliance:

not specified

The thermodynamic modelling indicates that when 2 μM of M.I. is added to seawater the major iodine species available for reaction with organic matter are HOI which increases with increasing pH, and 12 and I2C1- which decrease with increasing pH. HOI is a significant species over the entire range of 4-10 and until OI- predominates at very high pH values. The concentration of the hypoiodite anion, OI-, only becomes important at pH values above that of seawater; at pH 8 it represents only about 0.25% of the M.I added originally. At low pH values, I2 is likely to be the primary reactant because it is an electron acceptor and I2C1- is not; there HOI is not present in significant quantities.

Validity criteria fulfilled:

not specified

Conclusions:

HOI is a significant species over the entire range of 4-10 and until OI- predominates at very high pH values.

Executive summary:

The reactivity of 2 μM molecular iodine in seawater toward various organic compounds containing aromatic, a-keto, amino, olefinic and sugar functional groups was investigated. More detailed studies have been made of the reduction kinetics with salicylic acid,a-ketoglutaric acid and the polypeptide oxidized glutathione, particularly to establish whether variation over the pH range 4-9 would provide a similar reduction reactivity or "fingerprint" to that of molecular iodine added to natural seawater. The data indicates that compounds with only one functional group react with first order kinetics whereas compounds with multiple functional groups show more complex behaviour. Kinetic and thermodynamic modeling indicates that HOI is the main iodine species reacting with organic matter at seawater pH of 8.2. Based on the pH "fingerprints", peptides and compounds containing carbonyl or a-keto groups are the key reductants of molecular iodine added to seawater. Thesecompounds form C-I and N-I bonds which can allow for a rich organic iodine chemistry in seawater.

Reference 6

Endpoint:

hydrolysis

Type of information:

experimental study

Adequacy of study:

weight of evidence

Reliability:

2 (reliable with restrictions)

Rationale for reliability incl. deficiencies:

other: Study well documented, meets generally accepted scientific principles, acceptable for assessment in context of a weight of evidence approach

Qualifier:

no guideline followed

Principles of method if other than guideline:

In principle a spectrophotometric version of the classical iodine titration has been used. The iodine content was determined by measuring the absorbance at 285 nm.

GLP compliance:

no

Details on sampling:

For the disproportion reaction mixtures of 0.15 M sodium dihydrogenphosphate and 0.15 M sodium hydroxide were made up to 262 µmol/L by addition of saturated iodine solution in destilled water. By varying of the ratio of phosphate and hydroxide different pH values were archived.
0.4 mL of each reaction solution was mixed with 40 mL of 0.15 M sodium dihydrogenphosphate buffer solution (pH 6.7) and an excess of iodide (3.0 mL KI solution (0.6 M)). Afterwards the concentration of iodine was measured. All experiments were performed at 25 °C.

Buffers:

0.15 mol/L sodium dihydrogen phosphate buffer solution (pH 6.7) for determination of the iodine concentration, with an excess of iodide (3.0 mL of 0.60 M potassium iodide)

Details on test conditions:

Experiments were conducted under ordinary laboratory fluorescent lighting ; apart from the routine shielding of the potassium iodide solution from any light no other precautions were deemed necessary.

Transformation products:

not measured

Details on hydrolysis and appearance of transformation product(s):

The hydrolysis of iodine is a very complex system of competing reactions with two different equilibria in which various iodine species are formed and transformed in different pathways. The overall reaction can be summarized as 3 I2 + 3 H2O = IO3- + 5 I- + 6 H+. Hypoiodous acid (HOI) seems to be the key intermediate in the whole system.

pH:

7.99

Temp.:

25 °C

Hydrolysis rate constant:

0 min-1

DT50:

218 h

Type:

second order

Remarks on result:

other: calculated for PNECfreshwater

pH:

8.9

Temp.:

25 °C

Hydrolysis rate constant:

0.001 min-1

DT50:

52.2 h

Type:

second order

Remarks on result:

other: calculated for PNECfreshwater

pH:

10

Temp.:

25 °C

Hydrolysis rate constant:

0.001 min-1

DT50:

50.1 h

Type:

second order

Remarks on result:

other: calculated for PNECfreshwater

pH:

10.57

Temp.:

25 °C

Hydrolysis rate constant:

0 min-1

DT50:

164 h

Type:

second order

Remarks on result:

other: calculated for PNECfreshwater

pH:

11.1

Temp.:

25 °C

Hydrolysis rate constant:

0 min-1

DT50:

531 h

Type:

second order

Remarks on result:

other: calculated for PNECfreshwater

pH:

12.06

Temp.:

25 °C

Hydrolysis rate constant:

0 min-1

DT50:

573 d

Type:

second order

Remarks on result:

other: calculated for PNECfreshwater

Following gradients could be achieved by plotting the data of reciprocal concentration vs. time.

pH	no. of measuring points	gradient [L/mol*min]	correlation coefficient
12.06	30	23.2±0.2	0.9984
11.10	21	596±9	0.9962
10.57	25	1930±40	0.9943
10.00	26	6320±80	0.9981
8.90	25	6060±180	0.9898
7.99	27	1450±30	0.9960
7.18	28	-	-

Table 1

For pH 7.18 a second-order plot was not appropriate for fitting the experimental data. The experiments confirmed the observations of Thomas et al. (Thomas, 1980) that for pH 6 to 8 the solutions retain their brown colour for periods of days or weeks at 25 °C.

From the results of the performed experiments and former studies of the reaction state at 1 and 40 min, the maximum rate for the hydrolysis appears at ca. pH 9.5. At this point the species HOI predominates the reaction system and leads to the asssumption that hypoiodous acid plays a central role in the mechanism. Based on this assumption the reaction might be determined by a second-order reaction in [HOI], d([sum of iodine])/dt = -k*[HOI]².

The equilibria is described by

(i)                     I₂ +H₂O = HOI + H⁺ + I^-               K₁ = [HOI][H⁺][I^-]/[H₂O][I₂] (5.44E-13 mol/l, Allen et al., 1955)

(ii)                    HOI = H⁺ + OI^-                             K₂ = [H⁺][OI^-]/[HOI] (2.3E-11 mol/l, Palmer et al., 1982)

(iii)                   I₂ + I^- = I₃^-                                      K₃ = [I₃^-]/[I₂][I^-]       

By rearranging the equilibria equations and combination with the second-order equation following expression can be derived:

k_obs = k/(1 +(K₃[H⁺][I^-]²/K₁)+([H⁺][I^-]/K₁[H₂O])+K₂/[H⁺])²

Fitting the predicted values of this equation to the experimental ones can be achieved by trial and error alterations of the parameters [I^-] and k. A good correlation between the predicted and experimental results were obtained with k= 9E+3 L/mol and a iodide concentration of 1E-4 mol/L.

References:

Allen TL, Keefer RM (1955). The Formation of Hypoiodous Acid and Hydrated Iodine Cation by the Hydrolysis of Iodine, J Am Chem Soc, 77, 2957 -2960.

Palmer DA, Lietzke MH (1982). Radiochim Acta, 31, 37.

Thomas TR, Pence DT, Hasty RA (1980). The disproportionation of hypoiodous acid, J Inorg Nucl Chem, 42, 183 -186.

Validity criteria fulfilled:

not applicable

Remarks:

data presented in context of a weight of evidence

Conclusions:

Although the presented study is not conducted according a standard test guideline, it meets generally accepted scientific principles and is well documented.

Each study for itself does not fulfill the requirements of this endpoint, therefore a combined weight of evidence approach is applied. Further the application of the standard test guideline OECD 111 is considered to be not adequate due to the complex system of hydrolysis of iodine. The standard protocol tends to the identification of the hydrolysis products and the rate of hydrolysis as a function of the pH. As the hydrolysis of iodine is a two-step equilibrium reaction with numerous competing reactions and iodine species an adaptation of the test guideline would be appropriate. Based on available and evaluated literature data we tried to draw a complete picture of the behavior of molecular iodine in water.

In this case the first ultra-fast equilibrium reaction plays a decisive role for the concentration of molecular iodine in water which makes the use of special analytical techniques necessary. Eigen and Kustin analyzed the reaction in acidic media and from the stoichiometry it can be concluded that at environmental pH the reaction will be promoted due to the removal of the formed H+-ions. Thus, a fast and nearly complete transformation can be assumed. The findings of Thomas et al. and Truesdale et al. indicate that the second reaction is of second-order kinetics at pHs between 6 and 10 and therefore depending on the concentration of iodide and hypoiodous acid. Additionally, it is obvious from the observations that, in principle, the initial phase of the second-step disproportionation at environmental pH is fast. Then the gradient of the reaction rate flattens reaching for a steady-state at pHs between 6 and 8 which could remain for several weeks. At the end of the initial phase which lasts for about one hour depending on environmental pH the major partition of iodine is disproportionated to iodide and iodate.
Thus, the presented weight of evidence is considered to be adequate for assessing the environmental risk and describing the fate and behavior of molecular iodine in the aquatic compartment at different pH.

Executive summary:

Truesdale et al. examined the disproportionation of hypoiodous acid. They determined the rate constants between pH 7 and 12 for the hydrolysis of iodine and were able to compile a model to fit their results by rearranging the equilibria equitations and combining it with a second order kinetics. Hence, it can be assumed that HOI plays a central role in the mechanism of hydrolysis as the maximum rate constant can be measured at ca. pH 9.5 where HOI predominates the equilibria.

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 E-13, 25 °C and K= 0.49E-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).

References:

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.

Truesdale VW, Luther GW (1995). Molecular iodine reduction by natural and model organic substances in seawater, Aq. Geochem., 1, 89 -104

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