Converting to sodium hypochlorite alleviates some of the safety concerns associated with storing and using gaseous chlorine. However, it does not address the negative impact of chlorinated compounds on water and air quality. Because of it's non-selectivity, NaOCl will react rapidly with organics as well as H2S. This is a key point; most odor control chemical addition is done to raw sewage (collection system or plant headworks) where the organic load is relatively high. The aforementioned side reactions result in increased hypochlorite usage and the formation of chlorinated by-products which may be toxic or refractory to biological treatment. When NaOCl reacts with low molecular weight organics, VOCs (volatile organic compounds) can be produced, posing air quality concerns.

The use of H2O2 eliminates both chlorine gas safety concerns and chlorinated by-product formation. Under conditions found in municipal wastewater, H2O2 has less tendency to react with organics than NaOCl, resulting in more efficient oxidant use. Furthermore, any residual H2O2 remaining after oxidation of the H2S, decomposes to oxygen and water. This additional dissolved oxygen helps to maintain aerobic conditions, inhibiting further H2S generation.

Commercially, sodium hypochlorite is available in several solution concentrations, with 12.5% (w/w) being the most common for bulk usage. This product contains 1.25 lbs of available chlorine per gallon. Hydrogen peroxide is available in strengths from 35-70% (w/w), with 50% H2O2 being the most common for wastewater treatment applications.

In aqueous solution at pH 7 (typical pH for municipal wastewater), the sodium hypochlorite oxidation of H2S yields predominantly sulfate, as follows:

Therefore, 8.76 mg/l NaOCl (100%) are theoretically required to oxidize 1 mg/l of H2S. In actual practice, NaOCl:H2S weight ratios required for effective odor control range from 5:1 to as much as 15:1. Hydrogen peroxide oxidizes H2S predominantly to colloidal sulfur at pH 7. This reaction proceeds accordingly:

Therefore, 1 mg/l H2O2 (100%) is theoretically required to oxidize 1 mg/l H2S. In practice, H2O2:H2S weight ratios usually vary from 1:1 - 3:1. The key point of the above chemistry, is that significantly less H2O2 (compared to NaOCl) is required to oxidize the same amount of H2S.

When comparing NaOCl and H2O2 as alternatives to gaseous chlorine, chemical costs are obviously an important factor. The table above compares the relative chemical costs for NaOCl and H2O2 at a facility using 1,200 lbs/day of gaseous chlorine to treat 200 lbs/day of H2S.

The unit costs for Cl2 and NaOCl were based on information provided by municipal treatment facilities. The H2O2 cost is based on current list price in bulk. (1995 Prices) Oxidant:H2S weight ratios are averages based on experience at various treatment facilities. The comparison above demonstrates that H2O2 is a more cost effective chlorine alternative than NaOCl. In this case, chemical costs for H2O2 provide a 42.5% saving compared to NaOCl.

Another consideration when comparing H2O2 and NaOCl are the volume requirements of each chemical. From the above example, the volume of NaOCl solution needed on a daily basis is about 12 times larger than that of H2O2, to treat the same amount of H2S. Therefore, larger capital outlays will be required for NaOCl storage vessels, pumping systems, and piping.

Although both products experience some decomposition during storage, 50% hydrogen peroxide loses less than 1% per year when stored according to manufacturer's specifications. On the other hand, 12.5% NaOCl solution under normal storage conditions can lose 1-2% per month. The chlorine gas resulting from NaOCl decomposition is scrubbed (extra cost) or vented to the atmosphere, contributing to treatment plant air emissions.

Please note: This article also appeared in Volume 2 of H2O2 Update.

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