Dechlorination - Hydrogen Peroxide Safe Alternative
Chlorine continues to be used throughout industry as a bleaching agent, a disinfectant, and an oxidizer. As a result, it is present in many industrial and municipal wastewaters in concentrations ranging from a few ppb to 1% or more. Industry experience has verified that chlorine residuals, even at very low levels, are toxic to certain fish and other aquatic life. In some areas of the United States an upper limit of 0.1 mg/L chlorine residual already applies to wastewater discharges.
Dechlorination can be accomplished by several means, the most widely used being sulfur dioxide – either as a gas (SO2) or as a salt (e.g., sodium metabisulfite). However, this method has several shortcomings:
- As a corrosive compressed gas, SO2 presents safety and handling problems which require expensive storage and containment facilities
- As a salt, the bisulfite solids must be dissolved into solution before use and the solution often requires heating through the colder months to prevent freezing
- An excess of SO2 must be added to destroy all the available chlorine. Excess SO2 is itself toxic to aquatic life and exerts an oxygen demand on the receiving stream
- The process leaves behind sulfate salts which may degrade the effluent for recycling or subsequent processing
Hydrogen peroxide (H2O2) is a safe, convenient alternative for many dechlorination needs, especially those involving "free available chlorine" -- as opposed to "combined available chlorine."
When elemental chlorine is dissolved in water, an equilibrium is established between chlorine, hypochlorous acid, and hypochlorite ion (Cl2, HOCl and OCl-, respectively). The relative amount of each present depends primarily on the pH of the system.
|Cl2 + H2O ↔ HOCl + H+ + Cl-|
|HOCl ↔ OCl- + H+|
The chlorine as HOCl and OCl- is referred to as free available chlorine. This is the form of chlorine typically found in cooling water circuits, industrial bleaching systems, and many chemical processing operations. Nitrogen-containing compounds such as ammonia, amines and proteins are usually present in municipal wastewater. Free available chlorine reacts readily with these materials to form chloramines in which the chlorine is described as combined available chlorine. The available chlorine remaining after disinfection of municipal wastewaters is usually present in the combined form.
Hydrogen peroxide reacts with free available chlorine in solutions with pH > 7. While there is no upper limit to the pH (e.g., H2O2 can be used to dechlorinate effluent from caustic/chlorine odor scrubbers), as a practical matter, pH 8.5 is preferred in order to provide an instantaneous reaction.
About 0.48 pounds of hydrogen peroxide is required to destroy one pound of free available chlorine. In most cases the oxygen produced by the reaction will remain dissolved in the solution (saturation is about 10 ppm D.O.). Where higher concentrations of chlorine are involved, the solutions may effervesce and provision must be made to accommodate the O2 evolved. The reaction is mildly exothermic, liberating 37 kcal/mole as opposed to 199 kcal/mole when using SO2.
Significantly, hydrogen peroxide reacts very slowly with combined available chlorine. Consequently, solutions which contain ammonia (e.g., most municipal wastewater effluents) cannot be dechlorinated with H2O2.
Hydrogen peroxide should be investigated as a dechlorination agent in industrial waters characterized by free available chlorine. These include:
- Process effluent from chlorine production units and other chemical processes where chlorine is used as an oxidizing agent (including scrubbing of chlorine vapors).
- Cooling water blowdown where chlorination is used for microbiological control.
- Wash waters from industrial bleaching operations.
- Municipal wastewater effluent that has been denitrified prior to chlorination.
For example, hydrogen peroxide is being used to destroy chlorine in FMC’s wastewater from its chloralkali plant. Tests show 100% fish survival after 96 hours in the undiluted hydrogen peroxide-treated effluent. Other investigators report that hydrogen peroxide concentrations up to 40 mg/L have no effect on fingerling rainbow trout after 48 hours exposure (ref: Eden, Freske and Melbourne, Chemistry and Industry, p. 1105, Dec. 15, 1951). In fact, the dissolved oxygen from the hydrogen peroxide reaction with chlorine may improve the quality of the receiving water.
A simple test demonstrates the rapid destruction of free available chlorine by hydrogen peroxide. The free available chlorine concentration is determined in one liter of the test wastewater at a pH of 7 to 9. For each mg/L of free chlorine, add 1 mL of 0.1 wt.% H2O2. This quantity may be approximated by one or two drops of a 3 wt.% of hydrogen peroxide solution available in drug stores. The free chlorine will have disappeared by the time mixing is complete. The presence of the excess hydrogen peroxide used in this test does not interfere with the standard orthotolidine test for available chlorine.
For process control purposes, the simpler iodine titration may be used. Since the dechlorination reaction is quantitative and instantaneous, H2O2 and OCl- cannot coexist. Thus, although both reactants will liberate iodine from iodide, it is easy to adjust the analysis to determine which is present. If iodine is liberated in an alkali solution, the oxidant is hypochlorite which means more H2O2 should be added to complete the dechlorination. If no iodine is liberated in an alkali solution and if iodine is liberated in an acid solution (to which has been added ammonium molybdate) the oxidant is H2O2 which means that dechlorination is complete.
If little or no residual H2O2 is wanted in the dechlorinated effluent, it is possible to control the process by using a simple spot test to verify the presence or absence of H2O2. One such test involves placing onto a spot plate a few drops of 10 g/L CrO3 solution (pH adjusted to 0.5 - 1.0 with sulfuric acid), mixed with a few drops of the sample. If H2O2 is present, a rapidly fading blue color is obtained which corresponds to the formation of an unstable peroxo chromium complex. Hypochlorite cannot form this complex. A similar test can be done using TiCl3 which forms a yellow color which is specific to H2O2 (ref: Spot Tests in Inorganic Analysis, F. Feigl, Elsevier Publishing Corp., pg 529 (1958).
The above tests are more suited for controlling batch operations. For controlling continuous processes, online UV analysis may be employed which uses the differential absorptions of H2O2 (at 254 nm) and hypochlorite (at 313 nm). A two flow analyzer may be programmed to control the H2O2 metering pumps so that only mg/L levels of H2O2 are present in the reactor discharge. ORP (Oxidation Reduction Potential) can also be used as a process control tool provided that compensation is made for pH changes.
---FMC Technical Data, Pollution Control Release No. 5
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