Sulfur dioxide
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| Names | |||
|---|---|---|---|
| IUPAC name
Sulfur dioxide
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| Other names
Sulfurous anhydride
Sulfur(IV) ox | |||
| Identifiers | |||
3D model (JSmol)
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| 3535237 | |||
| ChEBI | |||
| ChEMBL | |||
| ChemSpider | |||
| ECHA InfoCard | 100.028.359 | ||
| EC Number |
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| E number | E220 (preservatives) | ||
| 1443 | |||
| KEGG | |||
| MeSH | Sulfur+dioxide | ||
PubChem CID
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| RTECS number |
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| UNII | |||
| UN number | 1079, 2037 | ||
CompTox Dashboard (EPA)
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| Properties | |||
| SO 2 | |||
| Molar mass | 64.066 g mol-1 | ||
| Appearance | Colorless gas | ||
| Density | 2.6288 kg m-3 | ||
| Melting point | −72 °C; −98 °F; 201 K | ||
| Boiling point | −10 °C (14 °F; 263 K) | ||
| 94 g dm-3[1] | |||
| Vapor pressure | 237.2 kPa | ||
| Acidity (pKa) | 1.81 | ||
| Basicity (pKb) | 12.19 | ||
| Viscosity | 0.403 cP (at 0 °C) | ||
| Structure | |||
| C2v | |||
| Digonal | |||
| Dihedral | |||
| 1.62 D | |||
| Thermochemistry | |||
Std molar
entropy (S⦵298) |
248.223 J K-1 mol-1 | ||
Std enthalpy of
formation (ΔfH⦵298) |
-296.81 kJ mol-1 | ||
| Hazards | |||
| NFPA 704 (fire diamond) | |||
| Lethal dose or concentration (LD, LC): | |||
LD50 (median dose)
|
3000 ppm (30 min inhaled, mouse) | ||
| Related compounds | |||
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).
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Sulfur dioxide (also sulphur dioxide) is the chemical compound with the formula SO
2. It is released by volcanoes and in various industrial processes. Since coal and petroleum often contain sulfur compounds, their combustion generates sulfur dioxide unless the sulfur compounds are removed before burning the fuel. Further oxidation of SO2, usually in the presence of a catalyst such as NO2, forms H2SO4, and thus acid rain.[2] Sulfur dioxide emissions are also a precursor to particulates in the atmosphere. Both of these impacts are cause for concern over the environmental impact of these fuels.
Structure and bonding
SO2 is a bent molecule with C2v symmetry point group. In terms of electron-counting formalism, the sulfur atom has an oxidation state of +4 and a formal charge of 0. It is surrounded by 5 electron pairs and can be described as a hypervalent molecule. From the perspective of molecular orbital theory, most of these valence electrons are engaged in S–O bonding.
Although sulfur and oxygen both have six valence electrons, the molecular bonds in SO2 are not the same as those in ozone. The S–O bonds are shorter in SO2 (143.1 pm) than in sulfur monoxide, SO (148.1 pm), whereas the O–O bonds are longer in ozone (127.8 pm) than in dioxygen, O2 (120.7 pm). The mean bond energy is greater in SO2 (548 kJ/mol) than in SO (524 kJ/mol), whereas it is less in O3 (297 kJ/mol) than in O2 (490 kJ/mol). These pieces of evidence lead chemists to conclude that the S–O bonds in sulfur dioxide have a bond order of at least 2, unlike the O–O bonds in ozone, which have a bond order of 1.5.[3]
Production
Combustion routes
Sulfur dioxide is the product of the burning of sulfur or of burning materials that contain sulfur:
- S8 + 8 O2 → 8 SO2
Sulfur dioxide is typically produced in significant amounts by the burning of common sulfur-rich materials including wool, hair, rubber, and foam rubber such as are found in mattresses, couch cushions, seat cushions, and carpet pads, and vehicle tires. Ferrous metals such as steel exposed to sulfur dioxide combustion fumes are rapidly oxidized and sulfidated. In house fires, this sometimes produces apparently molten steel comprising iron oxides and iron sulfide. The most common example of this phenomenon is “apparently melted steel” bedsprings that are found by fire investigators. The burning foam rubber in the mattress produces sulfur dioxide which reacts with the hot metal, further heating it until the oxide/sulfide melts, giving the appearance of “melted bed springs”. After the foam rubber burns away, a further heating of the apparently melted bedsprings in the presence of an excess of oxygen, re-releases the sulfur dioxide:
4 FeS2 + 11 O2 → 2 Fe2O3 + 8 SO2
The combustion of hydrogen sulfide and organosulfur compounds proceeds similarly.
- 2 H2S + 3 O2 → 2 H2O + 2 SO2
The roasting of sulfide ores such as pyrite, sphalerite, and cinnabar (mercury sulfide) also releases SO2:
- 4 FeS2 + 11 O2 → 2 Fe2O3 + 8 SO2
- 2 ZnS + 3 O2 → 2 ZnO + 2 SO2
- HgS + O2 → Hg + SO2
A combination of these reactions is responsible for the largest source of sulfur dioxide, volcanic eruptions. These events can release millions of tonnes of SO2.
Reduction of higher oxides
Sulfur dioxide is a by-product in the manufacture of calcium silicate cement: CaSO4 is heated with coke and sand in this process:
- 2 CaSO4 + 2 SiO2 + C → 2 CaSiO3 + 2 SO2 + CO2
The action of hot sulfuric acid on copper turnings produces sulfur dioxide.
- Cu + 2 H2SO4 → CuSO4 + SO2 + 2 H2O
From sulfite
Sulfite results from the reaction of aqueous base and sulfur dioxide. The reverse reaction involves acidification of sodium metabisulfite:
- H2SO4 + Na2S2O5 → 2 SO2 + Na2SO4 + H2O
Reactions
Industrial reactions
Treatment of basic solutions with sulfur dioxide affords sulfite salts:
- SO2 + 2 NaOH → Na2SO3 + H2O
Featuring sulfur in the +4 oxidation state, sulfur dioxide is a reducing agent. It is oxidized by halogens to give the sulfuryl halides, such as sulfuryl chloride:
- SO2 + Cl2 → SO2Cl2
Sulfur dioxide is the oxidising agent in the Claus process, which is conducted on a large scale in oil refineries. Here sulfur dioxide is reduced by hydrogen sulfide to give elemental sulfur:
- SO2 + 2 H2S → 3 S + 2 H2O
The sequential oxidation of sulfur dioxide followed by its hydration is used in the production of sulfuric acid.
- 2 SO2 + 2 H2O + O2 → 2 H2SO4
Laboratory reactions
Sulfur dioxide can react with certain 1,3-dienes in a cheletropic reaction to give organosulfur compounds.
Sulfur dioxide can bind to metal ions as a ligand to form metal sulfur dioxide complexes, typically where the transition metal is in oxidation state 0 or +1. Many different bonding modes (geometries) are recognized, but in most cases the ligand is monodentate, attached to the metal through sulfur, which can be either planar and pyramidal η1.[3]
Uses
Precursor to sulfuric acid
Sulfur dioxide is an intermediate in the production of sulfuric acid, being converted to sulfur trioxide, and then to oleum, which is made into sulfuric acid. Sulfur dioxide for this purpose is made when sulfur combines with oxygen. The method of converting sulfur dioxide to sulfuric acid is called the contact process. Several billion kilograms are produced annually for this purpose.
As a preservative
Sulfur dioxide is sometimes used as a preservative for dried apricots and other dried fruits owing to its antimicrobial properties, and it is sometimes called E220 when used in this way. As a preservative, it maintains the appearance of the fruit and prevents rotting.
In winemaking
Sulfur dioxide is an important compound in winemaking, and is designated as parts per million in wine, E number: E220.[4] It is present even in so-called unsulfurated wine at concentrations of up to 10 milligrams per litre.[5] It serves as an antibiotic and antioxidant, protecting wine from spoilage by bacteria and oxidation. Its antimicrobial action also helps to minimise volatile acidity. Sulfur dioxide is responsible for the words "contains sulfites" found on wine labels.
Sulfur dioxide exists in wine in free and bound forms, and the combination are referred to as total SO2. Binding, for instance to the carbonyl group of acetaldehyde, varies with the wine in question. The free form exists in equilibrium between molecular SO2 (as a dissolved gas) and bisulfite ion, which is in turn in equilibrium with sulfite ion. These equilibria depend on the pH of the wine. Lower pH shifts the equilibrium towards molecular (gaseous) SO2, which is the active form, while at higher SO2 more is found in the inactive sulfite and bisulfite forms. It is the molecular SO2 which is active as an antimicrobial and antioxidant, and this is also the form which may be perceived as a pungent odour at high levels. Wines with total SO2 concentrations below 10 ppm do not require "contains sulfites" on the label by US and EU laws. The upper limit of total SO2 allowed in wine in the US is 350 ppm; in the EU it is 160 ppm for red wines and 210 ppm for white and rosé wines. In low concentrations SO2 is mostly undetectable in wine, but at free SO2 concentrations over 50ppm, SO2 becomes evident in the nose and taste of wine.[citation needed]
SO2 is also a very important compound in winery sanitation. Wineries and equipment must be kept clean, and because bleach cannot be used in a winery[citation needed], a mixture of SO2, water, and citric acid is commonly used to clean and sanitize equipment. Compounds of ozone (O3) are now used extensively as cleaning products in wineries[citation needed] due to their efficiency, and because these compounds do not affect the wine or equipment.
As a reducing agent
Sulfur dioxide is also a good reductant. In the presence of water, sulfur dioxide is able to decolorize substances. Specifically it is a useful reducing bleach for papers and delicate materials such as clothes. This bleaching effect normally does not last very long. Oxygen in the atmosphere reoxidizes the reduced dyes, restoring the color. In municipal wastewater treatment sulfur dioxide is used to treat chlorinated wastewater prior to release. Sulfur dioxide reduces free and combined chlorine to chloride.[6]
Biochemical and biomedical roles
Sulfur dioxide is toxic in large amounts. It or its conjugate base bisulfite is produced biologically as an intermediate in both sulfate-reducing organisms and in sulfur oxidizing bacteria as well. The role of sulfur dioxide in mammalian biology is not yet well understood.[7] Sulfur dioxide blocks nerve signals from the pulmonary stretch receptors (PSR's) and abolishes the Hering–Breuer inflation reflex.
As a refrigerant
Being easily condensed and possessing a high heat of evaporation, sulfur dioxide is a candidate material for refrigerants. Prior to the development of freons, sulfur dioxide was used as a refrigerant in home refrigerators.
As a reagent and solvent in the laboratory
Sulfur dioxide is a versatile inert solvent that has been widely used for dissolving highly oxidizing salts. It is also used occasionally as a source of the sulfonyl group in organic synthesis. Treatment of aryl diazonium salts with sulfur dioxide and cuprous chloride affords the corresponding aryl sulfonyl chloride, for example:[8]
As an air pollutant
Sulfur dioxide is a noticeable component in the atmosphere, especially following volcanic eruptions.[9] According to the United States Environmental Protection Agency (EPA) (as presented by the 2002 World Almanac or in chart form[10]), the following amount of sulfur dioxide was released in the U.S. per year, measured in thousands of short tons:
- 1970 31,161
- 1980 25,905
- 1990 23,678
- 1996 18,859
- 1997 19,363
- 1998 19,491
- 1999 18,867
Sulfur dioxide is a major air pollutant and has significant impacts upon human health. In addition the concentration of sulfur dioxide in the atmosphere can influence the habitat suitability for plant communities as well as animal life.[11] Sulfur dioxide emissions are a precursor to acid rain and atmospheric particulates. Due largely to the US EPA’s Acid Rain Program, the U.S. has witnessed a 33 percent decrease in emissions between 1983 and 2002. This improvement resulted in part from flue gas desulfurization, a technology that enables SO2 to be chemically bound in power plants burning sulfur-containing coal or oil. In particular, calcium oxide (lime) reacts with sulfur dioxide to form calcium sulfite:
- CaO + SO2 → CaSO3
Aerobic oxidation of the CaSO3 gives CaSO4, anhydrite. Most gypsum sold in Europe comes from flue gas desulfurization.
Sulfur can be removed from coal during the burning process by using limestone as a bed material in Fluidized bed combustion.[12]
Sulfur can also be removed from fuels prior to burning the fuel. This prevents the formation of SO2 because there is no sulfur in the fuel from which SO2 can be formed. The Claus process is used in refineries to produce sulfur as a byproduct. The Stretford process has also been used to remove sulfur from fuel. Re-Dox processes using iron oxides can also be used, for example, Lo-Cat [13] or Sulferox.[14]
Fuel additives, such as calcium additives and magnesium oxide, are being used in gasoline and diesel engines in order to lower the emission of sulfur dioxide gases into the atmosphere.[15]
As of 2006, China is the world's largest sulfur dioxide polluter, with 2005 emissions estimated to be 25.49 million tons. This amount represents a 27% increase since 2000, and is roughly comparable with U.S. emissions in 1980.[16]
Safety
Inhaling sulfur dioxide is associated with increased respiratory symptoms and disease, difficulty in breathing, and premature death.[17] In 2008, the American Conference of Governmental Industrial Hygienists reduced the Short-term exposure limit from 5 ppm to 0.25 ppm. The OSHA PEL is currently set at 5 ppm (13 mg/m3) time weighted average. NIOSH has set the IDLH at 100 ppm.[18]
A 2011 systematic review concluded that exposure to sulfur dioxide is associated with preterm birth.[19]
- Food additive
In the United States, the Center for Science in the Public Interest lists the two food preservatives, sulfur dioxide and sodium bisulfite, as being safe for human consumption except for certain individuals who may be sensitive to it, especially in large amounts.[20]
See also
- Sulfur trioxide
- Sulfur-iodine cycle
- National Ambient Air Quality Standards
- Homer City Generating Station
References
- ^ Lide, David R., ed. (2006). CRC Handbook of Chemistry and Physics (87th ed.). Boca Raton, Florida: CRC Press. ISBN 0-8493-0487-3.
- ^ Holleman, Arnold Frederik; Wiberg, Egon (2001), Wiberg, Nils (ed.), Inorganic Chemistry, translated by Eagleson, Mary; Brewer, William, San Diego/Berlin: Academic Press/De Gruyter, ISBN 0-12-352651-5
- ^ a b Greenwood, Norman N.; Earnshaw, Alan (1997). Chemistry of the Elements (2nd ed.). Butterworth-Heinemann. p. 700. ISBN 978-0-08-037941-8.
- ^ Current EU approved additives and their E Numbers, The Food Standards Agency website.
- ^ Sulphites in wine, MoreThanOrganic.com.
- ^ Tchobanoglous, George (1979). Wastewater Engineering (3rd ed.). New York: McGraw Hill. ISBN 007041677X.
- ^ Liu, D. (2010). "Sulfur dioxide: a novel gaseous signal in the regulation of cardiovascular functions". Mini-Reviews in Medicinal Chemistry. 10: 1039–1045.
- ^ R. V. Hoffman (1990). "m-Trifluoromethylbenzenesulfonyl Chloride". Organic Syntheses; Collected Volumes, vol. 7, p. 508.
- ^ http://volcanoes.usgs.gov/hazards/gas/index.php
- ^ National Trends in Sulfur Dioxide Levels, United States Environmental Protection Agency.
- ^ C.Michael Hogan. 2010. Abiotic factor. Encyclopedia of Earth. eds Emily Monosson and C. Cleveland. National Council for Science and the Environment. Washington DC
- ^ Lindeburg, Michael R. (2006). Mechanical Engineering Reference Manual for the PE Exam. Belmont, C.A.: Professional Publications, Inc. pp. 27–3. ISBN 978-1-59126-049-3.
- ^ Lo-Cat Process
- ^ Sulferox
- ^ http://www.sfainternational.com/library/Marine%20Emissions%20Abatement.pdf, Page 6.
- ^ China has its worst spell of acid rain, United Press International.
- ^ [1]U.S. Environmental Protection Agency
- ^ "NIOSH Pocket Guide to Chemical Hazards".
- ^ "Air pollution and birth outcomes: a systematic review". Environ Int. 37: 498–516. 2011. doi:10.1016/j.envint.2010.10.009. PMID 21112090.
{{cite journal}}: Unknown parameter|authors=ignored (help) - ^ "Center for Science in the Public Interest - Chemical Cuisine". Retrieved March 17, 2010.
External links
- United States Environmental Protection Agency Sulfur Dioxide page
- International Chemical Safety Card 0074
- IARC Monographs. "Sulfur Dioxide and some Sulfites, Bisulfites and Metabisulfites" v54. 1992. p131.
- NIOSH Pocket Guide to Chemical Hazards
- Food Intolerance Network - Sulfite factsheet
- Sulfur Dioxide, Molecule of the Month