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Environmental fate & pathways

Hydrolysis

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Description of key information

Sodium percarbonate rapidly dissolves in water and dissociates into sodium ions, carbonate ions and hydrogen peroxide:
2 Na2CO3·3H2O2 → 4 Na+ + 2 CO32- + 3 H2O2

Key value for chemical safety assessment

Additional information

As sodium percarbonate rapidly dissolves in water, a short description of the stability of hydrogen peroxide and sodium carbonate is given, and the following text is copied from the OECD SIDS document for sodium percarbonate (OECD 2006), section 2.2.3:

Hydrogen peroxide

Hydrogen peroxide is a reactive substance in the presence of other substances, elements, radiation, materials or cells. Both biotic and abiotic degradation processes are important routes in removal of hydrogen peroxide in the environment:

2 H2O2 → 2 H2O + O2

Abiotic degradation of hydrogen peroxide is due to either reaction with itself (disproportionation),or reaction with transition metals, organic compounds capable of reacting with hydrogen peroxide, reaction with free radicals, heat or light (European Commission, 2003b). Hydrogen peroxide is normally a short-lived substance in the environment but half-lives vary greatly depending on the circumstances.

Sodium carbonate

Both sodium and inorganic carbonate have a wide natural occurrence (UNEP, 1995; OECD, 2003). The sodium concentration was reported for a total number of 75 rivers in North and South America, Africa, Asia, Europe and Oceania, with a 10th percentile of 1.5 mg/L, mean of 28 mg/L and 90th percentile of 68 mg/l (UNEP, 1995). Also the bicarbonate (HCO3-) concentration was reported for a total number of 77 rivers in North-America, South-America, Asia, Africa, Europe and Oceania. The 10th percentile, mean and 90th-percentile were 20, 106 and 195 mg/L, respectively. An emission of sodium carbonate to water will result in an increase in alkalinity and tendency to raise the pH value:

CO32- + H2O → HCO3- + OH-

HCO3- + H+ → CO2 + H2O

In water the carbonate ion will re-equilibrate until an equilibrium is established. The increase in pH depends on the buffer capacity of the water, which in most cases is determined by the natural background concentration of bicarbonate.