Chlorine dioxide: Difference between revisions
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==Further reading== |
==Further reading== |
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* [http://www.watersolutions.org.uk '''The Chlorine Dioxide Consultancy''' Whatever the Generation Technology] |
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* [http://www.clearwater.eu.com/chlorine_dioxide/how_it_works.html Catalytic Chlorine Dioxide Generation Method] |
* [http://www.clearwater.eu.com/chlorine_dioxide/how_it_works.html Catalytic Chlorine Dioxide Generation Method] |
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* [http://www.akzonobel.com/eka/products/pulp_and_paper_industry/plant_equipment_it/eka_engineering/products_services/index.aspx Major ClO2 supplier] |
* [http://www.akzonobel.com/eka/products/pulp_and_paper_industry/plant_equipment_it/eka_engineering/products_services/index.aspx Major ClO2 supplier] |
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Revision as of 07:29, 27 October 2009
 | This article needs additional citations for verification. (December 2007) |
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| Identifiers | |
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3D model (JSmol)
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| ChemSpider | |
| ECHA InfoCard | 100.030.135 |
| EC Number |
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PubChem CID
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| RTECS number |
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CompTox Dashboard (EPA)
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| Properties | |
| ClO2 | |
| Molar mass | 67.45 g/mol |
| Appearance | yellowish gas or liquid |
| Density | 3.04 g/cm3 |
| Melting point | -59.5 °C |
| Boiling point | 11 °C |
| 0.8 g/100 mL (20 °C) | |
| Solubility | soluble in alkalies, sulfuric acid |
| Acidity (pKa) | 2.5-3.5 |
| Thermochemistry | |
Std molar
entropy (S⦵298) |
257.22 J K−1 mol−1 |
Std enthalpy of
formation (ΔfH⦵298) |
+104.60 kJ/mol |
| Hazards | |
| Lethal dose or concentration (LD, LC): | |
LD50 (median dose)
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292 mg/kg (oral, rat) |
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).
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Chlorine dioxide is a chemical compound with the formula ClO2. This yellowish-green 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.[1]
Structure and bonding
The molecule ClO2 has an odd number of valence electrons and it is therefore a paramagnetic radical. Its electronic structure has baffled chemists for a long time because none of the possible Lewis structures is very satisfactory. In 1933 L.O. Brockway proposed a structure that involved a three-electron bond.[2] Pauling[3] further developed this idea and arrived at two resonance structures involving a double bond on one side and a single bond plus three-electron bond on the other. In Pauling's view the latter combination should represent a bond that is slightly weaker than the double bond. In molecular orbital theory this idea is commonplace if the third electron is placed in an anti-bonding orbital. Later work has confirmed that the HOMO is indeed an incompletely filled orbital.[citation needed]
Preparation
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:[4]
- 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 in a single reactor is believed to proceed 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.[5]
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.[citation needed] 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.
Uses
Chlorine dioxide is used primarily (>95%) for bleaching of wood pulp,[citation needed] but is also used for the bleaching of flour[citation needed] and for the disinfection of municipal drinking water.[citation needed] The Niagara Falls, New York water treatment plant first used chlorine dioxide for drinking water treatment in 1944 for phenol destruction.[citation needed] Chlorine dioxide was introduced as a drinking water disinfectant on a large scale in 1956,[citation needed] when Brussels, Belgium, changed from chlorine to chlorine dioxide.[citation needed] 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. [citation needed] Trihalomethanes are suspect carcinogenic disinfection by-products[citation needed] associated with chlorination of naturally occurring organics in the raw water.[citation needed] Chlorine dioxide is also superior to chlorine when operating above pH 7,[citation needed] in the presence of ammonia and amines[citation needed] and/or for the control of biofilms in water distribution systems.[citation needed] Chlorine dioxide is used in many industrial water treatment applications as a biocide including cooling towers, process water and food processing.[citation needed] Chlorine dioxide is less corrosive than chlorine and superior for the control of legionella bacteria.[citation needed]
It is more effective as a disinfectant than chlorine in most circumstances against water borne pathogenic microbes such as viruses[6] , 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.[citation needed] 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).[citation needed]
It can also be used for air disinfection,[citation needed] and was the principal agent used in the decontamination of buildings in the United States after the 2001 anthrax attacks.[citation needed] 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.[citation needed]
Chlorine dioxide is used as an oxidant for phenol destruction in waste water streams,[citation needed] control of zebra and quagga mussels in water intakes[citation needed] and for odor control in the air scrubbers of animal byproduct (rendering) plants.[citation needed]
Stabilized chlorine dioxide can also be used in an oral rinse to treat oral disease and malodor.[citation needed]
References
- ^ Greenwood, Norman N.; Earnshaw, Alan (1997). Chemistry of the Elements (2nd ed.). Butterworth-Heinemann. pp. 844–849. ISBN 978-0-08-037941-8.
- ^ Brockway LO (1933). "The Three-Electron Bond in Chlorine Dioxide". Proc. Natl. Acad. Sci. U.S.A. 19 (3): 303–7. PMC 1085967. PMID 16577512.
{{cite journal}}: Unknown parameter|month=ignored (help) - ^ Pauling, Linus (1988). General chemistry. Mineola, NY: Dover Publications, Inc. ISBN 0-486-65622-5.
- ^ Derby, R. I.; Hutchinson, W. S. "Chlorine(IV) Oxide" Inorganic Syntheses, 1953, IV, 152-158.
- ^ White, George W.; Geo Clifford White (1999). The handbook of chlorination and alternative disinfectants (4th ed.). New York: John Wiley. ISBN 0-471-29207-9.
{{cite book}}: CS1 maint: multiple names: authors list (link) - ^ Ogata N, Shibata T (2008). "Protective effect of low-concentration chlorine dioxide gas against influenza A virus infection". J. Gen. Virol. 89 (Pt 1): 60–7. doi:10.1099/vir.0.83393-0. PMID 18089729.
{{cite journal}}: Unknown parameter|month=ignored (help)