Chlorine dioxide

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Chlorine dioxide is a chemical compound with the formula ClO2. This reddish-yellow gas crystallizes as orange crystals at −59 °C. As one of several oxides of chlorine, it is a potent and useful oxidizing agent used in water treatment and in bleaching.


Chlorine dioxide is used primarily (>95%) for bleaching of wood pulp, but is also used for the bleaching of flour and for the disinfection of municipal drinking water. The Niagara Falls, New York water treatment plant first used chlorine dioxide for drinking water treatment in 1944 for phenol destruction. Chlorine dioxide was introduced as a drinking water disinfectant on a large scale in 1956, when Brussels, Belgium, changed from chlorine to chlorine dioxide. Its most common use in water treatment is as a pre-oxidant prior to chlorination of drinking water to destroy natural water impurities that produce trihalomethanes on exposure to free chlorine. Trihalomethanes are suspect carcinogenic disinfection by-product associated with chlorination of naturally occurring organics in the raw water. Chlorine dioxide is also superior to chlorine when operating above pH7, in the presence of ammonia and amines and/or for the control of biofilms in water distribution systems. Chlorine dioxide is used in many industrial water treatment applications as a biocide including cooling towers, process water and food processing. Chlorine dioxide is less corrosive than chlorine and superior for the control of legionella bacteria.

It is more effective as a disinfectant in most circumstances than chlorine against water borne pathogenic microbes such as viruses, bacteria and protozoa – including the cysts of Giardia and the oocysts of Cryptosporidium.

The use of chlorine dioxide in water treatment leads to the formation of the by-product chlorite which is currently limited to a maximum of 1 ppm in drinking water in the USA. This EPA standard limits the use of chlorine dioxide in the USA to relatively high quality water or water which is to be treated with iron based coagulants. (Iron can reduce chlorite to chloride.)

Protective effect of low-concentration chlorine dioxide gas against influenza A virus infection Ogata N, Shibata T. J Gen Virol 89 (2008), 60-67; DOI 10.1099/vir.0.83393-0

It can also be used for air disinfection, and was the principal agent used in the decontamination of buildings in the United States after the 2001 anthrax attacks. Recently, after the disaster of Hurricane Katrina in New Orleans, Louisiana and the surrounding Gulf Coast, chlorine dioxide has been used to eradicate dangerous mold from houses inundated by water from massive flooding.

Chlorine dioxide is used as an oxidant for phenol destruction in waste water streams, control of zebra and quagga mussels in water intakes and for odor control in the air scrubbers of animal byproduct (rendering) plants.

Stabilized chlorine dioxide can also be used in an oral rinse to treat oral disease and malodor, but its adverse side-effects are still being investigated.[1]


Chlorine dioxide is a highly endothermic compound that can decompose extremely violently when separated from diluting substances. As a result preparation methods that involve producing solutions of it without going through a gas phase stage are often preferred.

In the laboratory, ClO2 is prepared by oxidation of sodium chlorite:[2]

2NaClO2 + Cl2 → 2ClO2 + 2 NaCl

Over 95% of the chlorine dioxide produced in the world today is made from sodium chlorate and is used for pulp bleaching. It is produced with high efficiency by reducing sodium chlorate in a strong acid solution with a suitable reducing agent such as hydrochloric acid and sulfur dioxide. The reaction of sodium chlorate with hydrochloric acid proceeds in one reactor via the following pathway:

HClO3 + HCl → HClO2 + HOCl
HClO3 + HClO2 → 2ClO2 + Cl2 + 2H2O
HOCl + HCl → Cl2 + H2O

A much smaller but important market for chlorine dioxide is for use as a disinfectant. Since 1999 a growing proportion of the chlorine dioxide made globally for water treatment and other small scale applications has been made using the chlorate, hydrogen peroxide and sulfuric acid method which can produce a chlorine free product at high efficiency. Traditionally, chlorine dioxide for disinfection applications has been made by one of three methods using sodium chlorite or the sodium chlorite - hypochlorite method:

2NaClO2 + 2HCl + NaOCl → 2ClO2 + 3NaCl + H2O

or the sodium chlorite - hydrochloric acid method:

5NaClO2 + 4HCl → 5NaCl + 4ClO2 + 2H2O

All three sodium chlorite chemistries can produce chlorine dioxide with high chlorite conversion yield, but unlike the other processes the chlorite-HCl method produces completely chlorine free chlorine dioxide but suffers from the requirement of 25% more chlorite to produce an equivalent amount of chlorine dioxide.

Very pure chlorine dioxide can also be produced by electrolysis of a chlorite solution:

2NaClO2 + 2H2O   →   2ClO2 + 2NaOH + H2

High purity chlorine dioxide gas (7.7% in air or nitrogen) can be produced by the Gas:Solid method, which reacts dilute chlorine gas with solid sodium chlorite.

2NaClO2 + Cl2   →   2ClO2 + 2NaCl

These processes and several slight variations have been reviewed.[3]

Handling properties

At concentrations greater than 15% volume in air at STP, ClO2 explosively decomposes into chlorine and oxygen. The decomposition is initiated by light. Thus, it is never handled in concentrated form, but is almost always used as a dissolved gas in water in a concentration range of 0.5 to 10 grams per liter. Its solubility increases at lower temperatures: it is thus common to use chilled water (5 °C or 41 °F) when storing at concentrations above 3 grams per liter. In many countries, such as the USA, chlorine dioxide gas may not be transported at any concentration and is almost always produced at the application site using a chlorine dioxide generator. In some countries, chlorine dioxide solution below 3 grams per liter in concentration may be transported by land, but are relatively unstable and deteriorate quickly.

Stabilized chlorine dioxide

A number of products are marketed as "stabilized chlorine dioxide" (SCD). Most of these solutions do not actually contain chlorine dioxide but consist of solutions of buffered sodium chlorite. A weak acid can be added to SCD to "activate" it and make chlorine dioxide in-situ without a chlorine dioxide generator. However, a new formula for stabilized chlorine dioxide has recently been developed that is self-stable and used as a broad spectrum disinfectant and anti-microbial. This form of chlorine dioxide is currently being used against bacterial and viral outbreaks including MRSA, Legionella, and Norovirus. The use of SCD is effective when the demand for chlorine dioxide is low and when impurities, such as small amounts of sodium, can be tolerated. For application requiring above 5 kg day−1 ClO2, chlorine dioxide produced by a generator with either sodium chlorite or sodium chlorate is typically more economical.


  1. US 4689215 
  2. Derby, R. I.; Hutchinson, W. S. "Chlorine(IV) Oxide" Inorganic Syntheses, 1953, IV, 152-158.
  3. White, G. C. "Handbook of Chlorination and Alternative Disinfectants", 4th Edition (Wiley, 1999).

External links

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