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In an incident report related to an accident of a truck that was transporting thioglycolic acid, it is reported that almost 60 drums were spread out in a deep ravine near the town of Béziers (South France) on September 18, 2003 (Devaux, TGA truck accident, Report 143B12, 2003).

Twelve hours after the accident almost no thioglycolate was detected in samples taken between 2 and 30 km downstream of the accident in the small river close to the accident. Big amounts of dithiodiglycolate (up to 1000-2000 ppm; detection limit 0.4 ppm by ionic chromatography and 0.4 or 0.05 ppm by HPLC)) were measured. For 20 days, the amount of thioglycolate and dithiodiglycolate in 12 sampling points downstream of the accident was followed. Thioglycolate was never detected. The concentration of dithiodiglycolate decreased rapidly after the accident. At all sampling points the concentration is below 2 ppm three weeks after the accident. It must be noticed that 30 km downstream of the accident where drinking water is used from the river, neither thioglycolate nor dithiodiglycolate were detected during these 20 days. It should be noticed that no thioglycolate, but only dithiodiglycolate was detected in the soil at the place of the spillage.

Many experiments were conducted to understand what happened to thioglycolate in the environment.

Thioglycolate in water is only very slowly oxidized by air in acidic conditions, even in the presence of traces of iron (III) or manganese (II). However, thioglycolate in water, at a neutral pH, in the presence of 20 ppm of iron III or 20 ppm of manganese II salt can be oxidized by air to dithiodiglycolate within less than 3 hours. So, after the accident, thioglycolate has been neutralized by sediments containing calcium carbonate / magnesium salts then has been oxidized by catalytical amount of Fe (III) or Mn (II) contained in the ground. A very strong red/purple color that disappears after hours or days has been observed in the soil and water at the place of spillage. It has been demonstrated that it is caused by a thioglycolate-iron II complex at neutral pH. Once thioglycolate is oxidized to dithiodiglycolate, this color disappears. The extremely red violet color observed several days at the accident site is typical for the complex of thioglycolate with Fe (II) at nearly neutral pH. The color of this complex is so extreme, that a 1 ppm solution has such a high extinction, that it cannot be measured quantitatively by UV/VIS photometry with a normal 1 cm cuvette at 530 nm. The observed coloration occurs at concentrations below 1 ppm thioglycolate.

By reproducing in laboratory the contact of thioglycolate with the ground or the sediments of the small river, it was demonstrated that thioglycolate is quantitatively oxidized to dithiodiglycolate and that there is no other decomposition compound after 2 days (Devaux, 2003).

The amount of oxidized thioglycolic acid was determined in aerated tap water (Sablowski and Heitsch, 2004; reliability 2). The storage condition was an open beaker without stirring the solution. Samples were taken as fast as the HPLC-system could carry out one analytical run; that means approximately 35 min. The content of thioglycolate and of the oxidation product dithiodiglycolate were quantitatively determined. The result (See Figure 1,Oxidation of 100 ppm TGA in water, open beaker) described in this experiment is in good accordance with the truck accident here above cited, TGA half-live was approximately 35 min.

Some references provide answers related to oxidation kinetics of thioglycolate.

The pH and amount of oxidation occurring in thioglycolate solutions is influenced by the method of preparation of the solutions (Cook and Steel, 1958; reliability 2). For instance the effect of storage at different temperatures upon the oxidation of thioglycolate is not as great with unheated solutions as in heated solutions. Oxidation of thioglycolate is increased by dilution. The dithiodiglycolate produced on oxidation of thioglycolate itself, undergoes decomposition under alkaline conditions.

One percent thioglycolic acid solutions were prepared. Half of the solution was autoclaved at 115-116 °C for 30 minutes whilst the remainder was sterilized by passing through a 5/3 sintered glass filters. After sterilization, the solutions were stored in 100 mL glass-stoppered bottles at 20 °C in the dark. That thioglycolate content was determined by titration with potassium iodate solution in acid condition.

Thioglycolate content (% w/v) of nominal 1% solutions sterilized by autoclaving or filtration after storage at 20 °C was shown in the table below.

Time of storage in hours

0

24

90

120

Autoclaved

1.0

0.964

0.609

0.233

Filtered

0.907

0.871

0.677

0.534

The main loss of thioglycolic acid on storage is by oxidation to dithiodiglycolic acid. The oxidation is catalyzed by copper, manganese, iron and cobalt but not by zinc. The rate of oxidation varies with the pH, presence or absence of buffer and the concentration of metallic catalyst. The disulphide formed on oxidation also acts as a catalyst in the auto-oxidation of thioglycolic acid.

The effect of storage conditions upon thioglycolic acid are well shown by the appearance (change of color) of samples examined and their thioglycolic acid content after storage under different conditions (Cook and Steel, 1959; reliability 2). The older sample stored at a lower temperature had undergone less decomposition than a fresher sample kept at room temperature. These results are in agreement with the conclusions reached in Cook and Steel (1958), where the oxidation of thioglycolic acid was found to increase with dilution and temperature rise.

Bagiyan et al. (2003; reliability 2) has been studied self-oxidation of thiols. It was found that these reactions in neutral and alkaline solutions are induced by impurities of variable-valence metals. The ability of transition metals to catalyze oxidation rate versus pH passes through a maximum whose position on the pH scale depends on both the nature of metal and the structure of the thiol oxidized. The reactions of oxygen with thiol compounds in aqueous solutions are most often described by following equations:

2 RSH + O2-->RSSR + H2O2

4 RSH + O2-->2 RSSR + H2O

The reactions gave disulfides and hydrogen peroxide or water.

The thiol structure can affect the kinetics both directly (thiol can be involved in complexes limiting the reaction rate) and indirectly (by changing the concentration of the reactive thiolate form of thiol). Therefore, authors studied acid ionization of SH groups in solutions of thiol compounds having different structures by spectrometry in the UV range using absorption of the RS-chromophore in these compounds. According to the known mechanism, one would expect that the curve of self-oxidation rate Wo = f(pH) for thioglycolic acid will pass through a maximum at pH 10.5. In actuality, the maximum rates are attained at pH 6.5.

By the way, the limit of quantitation for thioglycolate with UV photometer of this complex is 0.02 ppm, this is in the level of graphite furnace atomic absorption spectrometry (a method for real trace analysis).

Literature search indicated that thioglycolic acid and its salts are not expected to undergo hydrolysis in the environment due to the lack of hydrolysable functional groups. Aqueous hydroxyl radical rate constants of 9.0 x 108, 3.6 x 109and 6.0 x 109L/mol*sec were determined for thioglycolic acid

at pH 1 (Buxton et al., 1988; reliability 4; Anbar and Neta, 1967; reliability 4; Dorfman and Adams, 1973; reliability 4), these values correspond to half-lives of 2.4 years, 220 and 130 days, respectively at an aqueous hydroxyl radical concentration 1.0 x 10-17mol/L (Mill et al., 1980; reliability 4).

In conclusion, the main process that leads to degradation of thioglycolic acid in water is a fast oxidation to dithiodiglycolate as demonstrated by literature, experiments and confirmed by the above mentioned incident report.

The thioglycolate is quantitatively oxidized to dithiodiglycolate, other decomposition compound after 2 days were not found. Further experiments demonstrated that dithiodiglycolate can undergo a rapid biodegradation at a concentration of 150 ppm, although some adaption of bacteria is necessary.