Chlorine dioxide

"ClO2" redirects here. For the oxoanion with the formula ClO−
2, see chlorite. For the oxycation with the formula ClO₂⁺, see Chloryl.

Chlorine dioxide
IUPAC name Chlorine dioxide
Other names Chlorine(IV) oxide Chloryl
Identifiers
CAS number	10049-04-4 Y
PubChem	24870
ChemSpider	23251 Y
UNII	8061YMS4RM Y
EC number	233-162-8
MeSH	Chlorine+dioxide
ChEBI	CHEBI:294​15 N
RTECS number	FO3000000
Gmelin Reference	1265
Jmol-3D images	Image 1 Image 2
SMILES O=[Cl]=O O=Cl[O]
InChI InChI=1S/ClO2/c2-1-3 Y Key: OSVXSBDYLRYLIG-UHFFFAOYSA-N Y InChI=1/ClO2/c2-1-3 Key: OSVXSBDYLRYLIG-UHFFFAOYAC
Properties
Molecular formula	ClO2
Molar mass	67.45 g mol−1
Exact mass	66.958681951 g mol-1
Appearance	Yellow gas
Odor	Acrid
Density	2.757 g dm−3[1]
Melting point	-59 °C, 214 K, -74 °F
Boiling point	11 °C, 284 K, 52 °F
Solubility in water	8 g dm-3 (at 20 °C)
Acidity (pKa)	3.0(5)
Thermochemistry
Std enthalpy of formation ΔfHo298	104.60 kJ mol-1
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)
N dioxide (verify) (what is: Y/N?) Except where noted otherwise, data are given for materials in their standard state (at 25 °C, 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

Pauling's proposal

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 are 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 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. 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 allows the best economy and does 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. Newer dosing systems have proven to provide a safer alternative producing ClO2 in solution and offering a huge leap in efficiency converting around 95% instantly only requiring the use of weak FDA approved acids offering a much safer method of producing ClO2 as there is no storage which proves a more effective Chlorine Dioxide Water Treatment.

Uses

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.^{[9][10]:4-1[11]} 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.^[12][13][14] Trihalomethanes are suspect carcinogenic disinfection by-products^[15] associated with chlorination of naturally occurring organics in the raw water.^[14] 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.^[14] Chlorine dioxide is used in many industrial water treatment applications as a biocide including cooling towers, process water and food processing.^[16]

Chlorine dioxide is less corrosive than chlorine and superior for the control of legionella bacteria.^[11][17] 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. Some unscrupulous biocide manufacturers will state that their product is EPA registered as a biocide. All EPA registered biocides must have a product label that is supplied with the product. This label will contain specifications as far as the product's EPA registration. EPA will register certain products as a general biocide, but others will have specifications for what bacteria the product can protect against. For instance, although chlorine dioxide is a registered biocide, it is not registered to protect against Legionella. If a biocide is sold without an EPA approved biocide label that is because the product is not registered as an EPA approved biocide.

It is more effective as a disinfectant than chlorine in most circumstances against water borne pathogenic microbes such as viruses,^[18] 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]}

It can also be used for air disinfection,^[19] and was the principal agent used in the decontamination of buildings in the United States after the 2001 anthrax attacks.^[20] 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.^[21]

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.^[10]:4-34

"Stabilized chlorine dioxide" is used in an oral rinse to treat oral disease and malodor.^[22] This term is a misnomer, referring simply to sodium chlorite, not chlorine dioxide.

Safety

On July 30, 2010 and again on October 1, 2010, the United States Food and Drug Administration 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 warning informed consumers that this industrial bleach can cause serious harm to health, including severe dehydration, nausea and diarrhea.^[23] There have been no reported deaths from use of MMS; however the FDA has received numerous reports of nausea, severe vomiting, and life-threatening low blood pressure caused by dehydration.^[24]

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.). Oxford: 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. doi:10.1073/pnas.19.3.303. PMC 1085967. PMID 16577512. http://www.pubmedcentral.nih.gov/articlerender.fcgi?tool=pmcentrez&artid=1085967. 
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.  edit
6. ^ Derby, R. I.; Hutchinson, W. S. "Chlorine(IV) Oxide" Inorganic Syntheses, 1953, IV, 152-158.
7. ^ ^{a b c d e} 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.  edit
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 0126785503. 
10. ^ ^{a b c d e f} EPA Guidance Manual, chapter 4: Chlorine dioxide, US Environmental Protection Agency, http://www.epa.gov/OGWDW/mdbp/pdf/alter/chapt_4.pdf, retrieved 2009-11-27 ^{[dead link]}
11. ^ ^{a b c d} Seymour Stanton Block (2001). Disinfection, sterilization, and preservation (5th ed.). Lippincott Williams & Wilkins. p. 215. ISBN 0683307401. 
12. ^ 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.  edit
13. ^ 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. 
14. ^ ^{a b c} 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. http://pubs.nrc-cnrc.gc.ca/rp/rppdf/s02-026.pdf. Retrieved 2009-11-27. 
15. ^ 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. http://www.pubmedcentral.nih.gov/articlerender.fcgi?tool=pmcentrez&artid=1569022. 
16. ^ 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.  edit
17. ^ 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). http://www.legionella.org/ZhangICHE07.pdf. Retrieved 2009-11-27. 
18. ^ 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. http://vir.sgmjournals.org/cgi/content/abstract/89/1/60?maxtoshow=&HITS=10&hits=10&RESULTFORMAT=&author1=ogata+n&andorexactfulltext=and&searchid=1&FIRSTINDEX=0&sortspec=relevance&resourcetype=HWCIT. 
19. ^ 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. http://www.csa.com/partners/viewrecord.php?requester=gs&collection=TRD&recid=07519213EN. 
20. ^ "Anthrax spore decontamination using chlorine dioxide". United States Environmental Protection Agency. 2007. http://www.epa.gov/opp00001/factsheets/chemicals/chlorinedioxidefactsheet.htm. Retrieved 2009-11-27. 
21. ^ 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. http://www.ingentaconnect.com/content/iafp/jfp/2005/00000068/00000006/art00007. 
22. ^ Frascella J.; Gilbert R. D.; Fernandez P.; Hendler J. (2000). "Efficacy of a chlorine dioxide-containing mouthrinse in oral malodor". Compend Contin Educ Dent 21 (3): 241–248. PMID 11199703. 
23. ^ FDA warning
24. ^ http://www.fda.gov/ForConsumers/ConsumerUpdates/ucm228052.htm