Chlorine Dioxide

Chlorine dioxide, discovered in 1811 by Davy, was prepared from the reaction of potassium chlorate with hydrochloric acid. Early experimentation showed that chlorine dioxide exhibited strong oxidizing and bleaching properties. In the 1930s, the Mathieson Alkali Works developed the first commercial process for preparing chlorine dioxide from sodium chlorate. By 1939, sodium chlorite was established as a commercial product for the generation of chlorine dioxide. Chlorine dioxide uses expanded rapidly in the industrial sector. In 1944, chlorine dioxide was first applied for taste and odor control at a water treatment plant in Niagara Falls, New York. Other water plants recognized the uses and benefits of chlorine dioxide. In 1958, a national survey determined that 56 U.S. water utilities were using chlorine dioxide. The number of plants using chlorine dioxide has grown more slowly since that time.

At present, chlorine dioxide is primarily used as a bleaching chemical in the pulp and paper industry. It is also used in large amounts by the textile industry, as well as for the &aching of flour, fats, oils, and waxes. In treating drinking water, chlorine dioxide is used in this country for taste and odor control, decolorization, disinfection, provision of residual disinfectant in water distribution systems, and oxidation of iron, manganese, and organics. The principal use of chlorine dioxide in the United States is for the removal of taste and odor caused by phenolic compounds in raw water supplies.

Chlorine dioxide is a yellow-green gas and soluble in water at room temperature to about 2.9 g/1 chlorine dioxide (at 30 mm mercury partial pressure) or more than 10 g/1 in chilled water. The boiling point of liquid chlorine dioxide is 11° C; the melting point is - 59° C. Chlorine dioxide gas has a specific gravity of 2.4. The oxidant is used in a water solution and is five times more soluble in water than chlorine gas. In addition, chlorine dioxide does not react with water in the same manner that chlorine does. Chlorine dioxide is volatile; consequently, it can be stripped easily from a water solution by aeration.

Chlorine dioxide has a disagreeable odor, similar to that of chlorine gas, and is detectable at 17 ppm. It is distinctly irritating to the respiratory tract at a concentration of 45 ppm in air. Concentrations above 11 percent can be mildly explosive in air. As a gas or liquid, it readily decomposes upon exposure to ultraviolet light. It is also sensitive to temperature and pressure, two reasons why chlorine dioxide is generally not shipped in bulk concentrated quantities. Chlorine dioxide has a much greater oxidative capacity than chlorine and is therefore a more effective oxidant in lower concentrations. Chlorine dioxide also maintains an active residual in potable water longer than chlorine does. It does not react with ammonia or with trihalomethane precursors when prepared with no free residual chlorine. Chlorine dioxide is prepared from feedstock chemicals by several methods. The specific method depends on the quantity needed and the safety limitations in handling the various feedstock chemicals. The most common processes are: From sodium chlorite (NaCIOJ:

• Acid and sodium chlorite

• Gaseous chlorine and sodium chlorite

• Sodium hypochlorite, acid, and sodium chlorite.

From sodium chlorate (NaCI03):

• The sulfur dioxide process

• The methanol process.

The first group of processes is more commonly used. The second group of processes is frequently used by industry where the quantities produced are much greater than in water utilities.

Oxidation of phenols with chlorine dioxide or chlorine produces chlorinated aromatic intermediates before ring rupture. Oxidation of phenols with either chlorine dioxide or ozone produces oxidized aromatic compounds as intermediates which undergo ring rupture upon treatment with more oxidant and/or longer reaction times. In many cases, the same nonchlorinated, ringruptured aliphatic products are produced using ozone or chlorine dioxide. In oxidizing organic materials, chlorine dioxide can revert back to the chlorite ion. In the presence of excess chlorine (or other strong oxidant), chlorite can be preoxidized to chlorine dioxide. Using large excesses of chlorine dioxide over the organic materials appears to favor oxidation reactions (without chlorination), but slight excesses appear to favor chlorination. When excess free chlorine is present with the chlorine dioxide, chlorinated organics usually are produced, but in lower yields, depending on the concentration of chlorine and its reactivity with the particular organic(s) involved. Treatment of organic compounds with pure chlorine dioxide containing no excess free chlorine produces oxidation products containing no chlorine in some cases, but products containing chlorine in others.

Under drinking water plant treatment conditions, humic materials and/ or resorcinol do not produce trihalomethanes with chlorine dioxide even when a slight excess of chlorine (1 percent to 2 percent) is present. Also, saturated aliphatic compounds are not reactive with chlorine dioxide. Alcohols are oxidized to the corresponding acids.

The gaseous chlorine-sodium chlorite process for producing chlorine dioxide uses aqueous chlorine and aqueous sodium chlorite to produce a mixture of chlorine dioxide and chlorine (commonly as HOC1). Figure 4 shows such a system, consisting of a chlorine dioxide generator, a gas chlorinator, a storage reservoir for liquid sodium chlorite, and a chemical metering pump. (Sodium chlorite solution can be prepared from commercially available dry chemical by adding it to water.) The recommended feed ratio of chlorine to sodium chlorite is 1:1 by weight. Additional chlorine can be injected into the reactor vessel without changing the overall production of chlorine dioxide.

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