Chlorine dioxide

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Chlorine dioxide
IUPAC name

Chlorine dioxide

Other names

Chlorine(IV) oxide
Chloryl

Identifiers
CAS number 10049-04-4 YesY
PubChem 24870
ChemSpider 23251 YesY
UNII 8061YMS4RM YesY
EC number 233-162-8
MeSH Chlorine+dioxide
ChEBI CHEBI:29415 YesY
RTECS number FO3000000
Gmelin Reference 1265
Jmol-3D images {{#if:O=[Cl]=OO=Cl[O]|Image 1
Image 2
Properties
Molecular formula ClO2
Molar mass 67.45 g mol−1
Appearance Yellow to reddish gas
Odor Acrid
Density 2.757 g dm−3[1]
Melting point −59 °C; −74 °F; 214 K
Boiling point 11 °C; 52 °F; 284 K
Solubility in water 8 g dm−3 (at 20 °C)
Solubility soluble in alkaline and sulfuric acid solutions
kH 4.01 x 10−2 atm-cu m/mole
Acidity (pKa) 3.0(5)
Thermochemistry
Std enthalpy of
formation ΔfHo298		 104.60 kJ/mol
Standard molar
entropy So298		 257.22 J K−1 mol−1
Hazards
MSDS ICSC 0127
EU Index 017-026-00-3
EU classification O T+ C N
R-phrases R6, R8, R26, R34, R50
S-phrases (S1/2), S23, S26, S28, S36/37/39, S38, S45, S61
NFPA 704
0
3
4
OX
LD50 292 mg/kg (oral, rat)
 YesY (verify) (what is: YesY/N?)
Except where noted otherwise, data are given for materials in their standard state (at 25 °C (77 °F), 100 kPa)
Infobox references

Chlorine dioxide is a chemical compound with the formula ClO2. This yellowish-green gas crystallizes as bright 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.[2]

Structure and bonding

Comparison of three-electron bond to the conventional covalent bond.
The two resonance structures

The molecule ClO2 has an odd number of valence electrons and therefore, it is a paramagnetic radical. Its electronic structure has long baffled chemists because none of the possible Lewis structures is very satisfactory. In 1933 L.O. Brockway proposed a structure that involved a three-electron bond.[3] Chemist Linus Pauling 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.[4] 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.[5]

Preparation

Chlorine dioxide is a 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. Arranging handling in a safe manner is essential.

In the laboratory, ClO2 is prepared by oxidation of sodium chlorite:[6]

2 NaClO2 + Cl2 → 2 ClO2 + 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 methanol, hydrogen peroxide, hydrochloric acid or sulfur dioxide.[7] Modern technologies are based on methanol or hydrogen peroxide, as these chemistries allow the best economy and do not co-produce elemental chlorine. The overall reaction can be written;

Chlorate + Acid + reducing agent → Chlorine Dioxide + By-products

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

The commercially more important production route uses methanol as the reducing agent and sulfuric acid for the acidity. Two advantages by not using the chloride-based processes are that there is no formation of elemental chlorine, and that sodium sulfate, a valuable chemical for the pulp mill, is a side-product. These methanol-based processes provide high efficiency and can be made very safe.[7]

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:

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

or the sodium chlorite - hydrochloric acid method:

5 NaClO2 + 4 HCl → 5 NaCl + 4 ClO2 + 2 H2O

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. Alternatively, hydrogen peroxide may efficiently be used also in small scale applications.[7]

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

2 NaClO2 + 2 H2O → 2 ClO2 + 2 NaOH + 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.

2 NaClO2 + Cl2 → 2 ClO2 + 2 NaCl

These processes and several slight variations have been reviewed.[8]

Handling properties

At gas phase concentrations greater than 30% volume in air at STP (more correctly: at partial pressures above 10 kPa [7]), ClO2 may explosively decompose into chlorine and oxygen. The decomposition can be initiated by, for example, light, hot spots, chemical reaction, or pressure shock. Thus, chlorine dioxide gas is never handled in concentrated form, but is almost always handled 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.[7] 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. The use of chlorine dioxide generators is steadily falling out of fashion in industry, as these systems generally require the use of strong acids to work and can take several hours to reach their full yield with poor efficiency. The requirement to store the hazardous gas in a pressurized chamber poses a risk some sites prefer not to take.

Uses

Chlorine dioxide is used primarily (>95%) for bleaching of wood pulp, and for the disinfection (called chlorination) of municipal drinking water.[9][10]:4-1[11]

Bleaching

Chlorine dioxide is sometimes used for bleaching of wood pulp in combination with chlorine, but it is used alone in ECF (elemental chlorine-free) bleaching sequences. It is used at moderately acidic pH (3.5 to 6). The use of chlorine dioxide minimizes the amount of organochlorine compounds produced.[12] Chlorine dioxide (ECF technology) currently is the most important bleaching method world wide. About 95% of all bleached Kraft pulp is made using chlorine dioxide in ECF bleaching sequences.[13]

Chlorine dioxide is also used for the bleaching of flour.

Water chlorination

Further information: Water chlorination

The Niagara Falls, New York water treatment plant first used chlorine dioxide for drinking water treatment in 1944 for phenol destruction.[10]:4–17[11] Chlorine dioxide was introduced as a drinking water disinfectant on a large scale in 1956, when Brussels, Belgium, changed from chlorine to chlorine dioxide.[11] 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.[14][15][16] Trihalomethanes are suspect carcinogenic disinfection by-products[17] associated with chlorination of naturally occurring organics in the raw water.[16] Chlorine dioxide is also superior to chlorine when operating above pH 7,[10]:4–33 in the presence of ammonia and amines[citation needed] and/or for the control of biofilms in water distribution systems.[16] Chlorine dioxide is used in many industrial water treatment applications as a biocide including cooling towers, process water, and food processing.[18]

Chlorine dioxide is less corrosive than chlorine and superior for the control of legionella bacteria.[11][19] Chlorine dioxide is superior to some other secondary water disinfection methods in that chlorine dioxide: 1) is an EPA registered biocide, 2) is not negatively impacted by pH 3) does not lose efficacy over time (the bacteria will not grow resistant to it) and 4) is not negatively impacted by silica and phosphate, which are commonly used potable water corrosion inhibitors.

It is more effective as a disinfectant than chlorine in most circumstances against water borne pathogenic microbes such as viruses,[20] bacteria and protozoa – including the cysts of Giardia and the oocysts of Cryptosporidium.[10]:4-20–4-21

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.[10]:4–33 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]

Other disinfection uses

It can also be used for air disinfection,[21] and was the principal agent used in the decontamination of buildings in the United States after the 2001 anthrax attacks.[22] 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 the flood-water.[23] Sometimes it is used as a fumigant treatment to 'sanitize' fruits such as blueberries, raspberries, and strawberries that develop molds and yeast.

Chlorine dioxide is used for the disinfection of endoscopes, such as, under the trade name Tristel.[24] It is also available in a "trio" consisting of a preceding "pre-clean" with surfactant and a succeeding "rinse" with deionised water and low-level antioxidant.[25]

Chlorine dioxide also is used for control of zebra and quagga mussels in water intakes.[10]:4–34

Chlorine dioxide also was shown to be effective in bedbug eradication.[26]

Other uses

Chlorine dioxide is used as an oxidant for phenol destruction in waste water streams and for odor control in the air scrubbers of animal byproduct (rendering) plants.[10]:4–34

Safety issues in water and supplements

Chlorine dioxide is toxic, hence limits on exposure to it are needed to ensure its safe use. The United States Environmental Protection Agency has set a maximum level of 0.8 mg/L for chlorine dioxide in drinking water.[27] The Occupational Safety and Health Administration (OSHA), an agency of the United States Department of Labor, has set a 8 hour permissible exposure limit of 0.1 ppm in air (0.3 milligrams per cubic meter (mg/m(3))) for people working with chlorine dioxide.[28]

On July 30, 2010 and again on October 1, 2010, the United States Food and Drug Administration (FDA), warned against the use of the product "Miracle Mineral Supplement" or "MMS", which when made up according to instructions produces chlorine dioxide. MMS has been marketed as a treatment for a variety of conditions, including HIV, cancer, and acne. The FDA warnings informed consumers that MMS can cause serious harm to health, and stated that it has received numerous reports of nausea, severe vomiting, and life-threatening low blood pressure caused by dehydration,[29][30] among other symptoms, such as diarrhea.

References

  1. ↑ Haynes, William M. (2010). Handbook of Chemistry and Physics (91 ed.). Boca Raton, Florida: CRC Press. p. 4–58. ISBN 978-1439820773. 
  2. ↑ Greenwood, Norman N.; Earnshaw, Alan (1997). Chemistry of the Elements (2nd ed.). Butterworth-Heinemann. pp. 844–849. ISBN 0080379419. 
  3. ↑ Brockway LO (March 1933). "The Three-Electron Bond in Chlorine Dioxide". Proc. Natl. Acad. Sci. U.S.A. 19 (3): 303–7. Bibcode:1933PNAS...19..303B. doi:10.1073/pnas.19.3.303. PMC 1085967. PMID 16577512. 
  4. ↑ Pauling, Linus (1988). General chemistry. Mineola, NY: Dover Publications, Inc. ISBN 0-486-65622-5. 
  5. ↑ Flesch, R.; Plenge, J.; Rühl, E. (2006). "Core-level excitation and fragmentation of chlorine dioxide". International Journal of Mass Spectrometry. 249-250: 68–76. Bibcode:2006IJMSp.249...68F. doi:10.1016/j.ijms.2005.12.046. 
  6. ↑ Derby, R. I.; Hutchinson, W. S. "Chlorine(IV) Oxide" Inorganic Syntheses, 1953, IV, 152-158.
  7. ↑ 7.0 7.1 7.2 7.3 7.4 Vogt, H.; Balej, J.; Bennett, J. E.; Wintzer, P.; Sheikh, S. A.; Gallone, P.; Vasudevan, S.; Pelin, K. (2010). "Chlorine Oxides and Chlorine Oxygen Acids". Ullmann's Encyclopedia of Industrial Chemistry. Wiley-VCH. doi:10.1002/14356007.a06_483.pub2. 
  8. ↑ 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. 
  9. ↑ Thomas Wilson Swaddle (1997). Inorganic chemistry: an industrial and environmental perspective. Academic Press. pp. 198–199. ISBN 0-12-678550-3. 
  10. ↑ 10.0 10.1 10.2 10.3 10.4 10.5 10.6 EPA Guidance Manual, chapter 4: Chlorine dioxide, US Environmental Protection Agency, retrieved 2009-11-27 
  11. ↑ 11.0 11.1 11.2 11.3 Seymour Stanton Block (2001). Disinfection, sterilization, and preservation (5th ed.). Lippincott Williams & Wilkins. p. 215. ISBN 0-683-30740-1. 
  12. ↑ E. Sjöström (1993). Wood Chemistry: Fundamentals and Applications. Academic Press. ISBN 0-12-647480-X. OCLC 58509724. 
  13. ↑ http://www.aet.org/science_of_ecf/eco_risk/2005_pulp.html
  14. ↑ Sorlini, S.; Collivignarelli, C. (2005). "Trihalomethane formation during chemical oxidation with chlorine, chlorine dioxide and ozone of ten Italian natural waters". Desalination 176 (1–3): 103–111. doi:10.1016/j.desal.2004.10.022. 
  15. ↑ Li J.; Yu Z.; Gao M. (1996). "A pilot study on trihalomethane formation in water treated by chlorine dioxide (translated from Chinese)". Zhonghua Yu Fang Yi Xue Za Zhi (Chinese journal of preventive medicine) 30 (1): 10–13. PMID 8758861. 
  16. ↑ 16.0 16.1 16.2 C. J. Volk; R. Hofmann; C. Chauret; G. A. Gagnon; G. Ranger; R. C. Andrews (2002). "Implementation of chlorine dioxide disinfection: Effects of the treatment change on drinking water quality in a full-scale distribution system". J. Environ. Eng. Sci. 1: 323–330. doi:10.1139/SO2-026. Retrieved 2009-11-27. 
  17. ↑ M. A. Pereira; L. H. Lin; J. M. Lippitt; S. L. Herren (1982). "Trihalomethanes as initiators and promoters of carcinogenesis". Environ Health Perspect 46: 151–156. doi:10.2307/3429432. JSTOR 3429432. PMC 1569022. PMID 7151756. 
  18. ↑ Andrews, L.; Key, A.; Martin, R.; Grodner, R.; Park, D. (2002). "Chlorine dioxide wash of shrimp and crawfish an alternative to aqueous chlorine". Food Microbiology 19 (4): 261–267. doi:10.1006/fmic.2002.0493. 
  19. ↑ Zhe Zhang; Carole McCann; Janet E. Stout; Steve Piesczynski; Robert Hawks; Radisav Vidic; Victor L. Yu (2007). "Safety and Efficacy of Chlorine Dioxide for Legionella control in a Hospital Water System". Infection Control and Hospital Epidemiology 28 (8). Retrieved 2009-11-27. 
  20. ↑ Ogata N, Shibata T (January 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. 
  21. ↑ Zhang, Y. L.; Zheng, S. Y.; Zhi, Q. (2007). "Air Disinfection with Chlorine Dioxide in Saps". Journal of Environment and Health 24 (4): 245–246. 
  22. ↑ "Anthrax spore decontamination using chlorine dioxide". United States Environmental Protection Agency. 2007. Retrieved 2009-11-27. 
  23. ↑ Sy, Kaye V.; McWatters, Kay H.; Beuchat, Larry R. (2005). "Efficacy of Gaseous Chlorine Dioxide as a Sanitizer for Killing Salmonella, Yeasts, and Molds on Blueberries, Strawberries, and Raspberries". Journal of Food Protection (International Association for Food Protection) 68 (6): 1165–1175. PMID 15954703. 
  24. ↑ Coates, D. (2001). "An evaluation of the use of chlorine dioxide (Tristel One-Shot) in an automated washer/disinfector (Medivator) fitted with a chlorine dioxide generator for decontamination of flexible endoscopes". Journal of Hospital Infection 48 (1): 55–65. doi:10.1053/jhin.2001.0956. PMID 11358471. 
  25. ↑ Tristel Wipes System Product Information at Ethical Agents, retrieved Nov 2012
  26. ↑ Gibbs SG, Lowe JJ, Smith PW, Hewlett AL., March 2012, Gaseous chlorine dioxide as an alternative for bedbug control, Department of Environmental, Agricultural and Occupational Health, University of Nebraska Medical Center College of Public Health
  27. ↑ "ATSDR: ToxFAQs™ for Chlorine Dioxide and Chlorite". 
  28. ↑ "Occupational Safety and Health Guideline for Chlorine Dioxide". Retrieved Dec 8, 1012. 
  29. ↑ FDA warning
  30. ↑ http://www.fda.gov/ForConsumers/ConsumerUpdates/ucm228052.htm

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