Sulfide Oxidation with Hydrogen Peroxide

Sulfide Odor Control

Sulfide is found throughout the environment as a result of both natural and industrial processes. Most sulfide found in nature was produced biologically (under anaerobic conditions) and occurs as free hydrogen sulfide (H2S) – characterized by its rotten egg odor. We are most likely to encounter biogenic H2S in sour groundwaters, swamps and marshes, natural gas deposits, and sewage collection/treatment systems. Manmade sources of H2S typically occur as a result of natural materials containing sulfur (e.g., coal, gas and oil) being refined into industrial products. For a variety of reasons – aesthetics (odor control), health (toxicity), ecological (oxygen depletion in receiving waters), and economic (corrosion of equipment and infrastructure) – sulfide laden wastewaters must be handled carefully and remediated before they can be released to the environment. Typical discharge limits for sulfide are < 1 mg/L.

Sulfide Treatment Alternatives

There are dozens of alternatives for treating sulfide laden waters, ranging from simple air stripping (for the low levels present in groundwaters) to elaborate sulfur recovery plants (used to treat several tons per day at refineries and coal burning power plants). There are processes based on biology (using compost filters, scrubbing media, or inhibition/disinfection), chemistry (oxidation, precipitation, absorption, and combination), and physics (adsorption, volatilization, and incineration). Each process occupies a niche which is often defined by the scale and continuity of treatment, whether the sulfide is in solution or is a gas, the concentration of sulfide involved, and the disposition of the sulfide containing medium. However, for reasons relating to convenience and flexibility, chemical oxidation (using hydrogen peroxide) continues to grow in its scope of application.

Odor Control Treatment with Hydrogen Peroxide

While other peroxygens such as permonosulfuric (Caro’s) acid, peracetic acid, and persulfates will oxidize sulfide, their use for this application is overkill. Hydrogen peroxide (H2O2) is considerably simpler and more cost-effective.

H2O2 may control sulfides in two ways, depending on the application:

• Prevention – by providing dissolved oxygen which inhibits the septic conditions which lead to biological sulfide formation; and
• Destruction – by oxidizing sulfide to elemental sulfur or sulfate ion.

This article focuses on the oxidation chemistry of odor control with H2O2, particularly as it is applied to industrial wastewaters containing moderate to high levels of sulfide (50 – 10,000 mg/L). Oxidation of sulfide with H2O2 proceeds differently depending primarily on the pH of the wastewater.

Neutral – Slightly Acid Conditions

The product of the oxidation is predominately elemental sulfur, which appears as a yellow colloid (if underdosed) or a white colloid (with complete oxidation). If clarity of the effluent is needed, the sulfur may be removed by flocculation with an anionic polymer followed by filtration.

The stoichiometry calls for 1.0 lb H2O2 per lb H2S, and it is not unusual for efficiencies to approach 100%, particularly when the concentrations of other oxidizable substances (e.g., thiosulfate) are low, and when the reaction is accelerated by catalysis. The effect of catalysis on the speed of reaction is illustrated in the following table:

Note: H2O2 dosed at 1.2:1 wt.% on sulfide, pH 7.0, and 20 deg-C in demineralized water.

There is very little heat generated in the reaction, even when sulfide levels are several thousand mg/L. Problems with heat and effervescence could arise however if the H2O2 was grossly overdosed in the presence of catalyst.

Alkaline Conditions

The above reaction predominates at pH > 9.2, and yields soluble sulfate as the reaction product. The stoichiometry calls for 4.25 lbs H2O2 per lb S 2-, and again it is not unusual for reaction efficiencies to approach 100%, provided that the H2O2 is added in a controlled fashion and the reaction medium is thoroughly mixed. This is due to the much faster reaction brought about by the increased reactivity of H2O2 at alkaline pH. Consequently, as the pH increases above 9-10, there is generally little benefit to catalyzing the reaction since, as the table below illustrates.

Note: H2O2 dosed at 5.0:1 wt.% on sulfide and 20 deg-C in demineralized water.

The theoretical heat of reaction is 225 kcal/gm-mol of SO4 2- (or 12,700 B.t.u. evolved per lb-S). Consequently, there may be a need to dissipate heat when treating solutions containing several thousand mg/L sulfide. Contributing to heat evolution is the unproductive decomposition of H2O2 which can become significant if H2O2 is overdosed (or improperly mixed), especially at pH > 11-12. The heat of H2O2 decomposition is about 1000 B.t.u. per lb-H2O2. H2O2 decomposition into oxygen, water and heat is a concern which is made more significant if the sulfide containing waste also contains volatile hydrocarbons.

Slightly Alkaline Conditions

In moving from pH 7 to pH 9, both of the above reactions may occur with the following results:

• The reaction products transition from elemental sulfur to sulfate
• The H2O2 requirement transitions from 1:1 to 4.25:1
• The rate of reaction speeds up

To some extent, catalysts may be used to push the reaction one way or the other. Catalysts such as iron, copper and manganese tend to favor sulfate formation, while those based on nickel and vanadium tend to favor elemental sulfur. These catalysts may be used to economize H2O2 use or to produce a clear effluent. In both cases, the speed of reaction is greatly accelerated, and fixed bed reaction columns (using e.g., zeolite catalyst support systems) may be used to lessen the environmental impact.

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