Sulfur dioxide

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Sulfur dioxide
Identifiers
CAS number7446-09-5 YesY
PubChem1119
ChemSpider1087 YesY
UNII0UZA3422Q4 YesY
EC number231-195-2
UN number1079, 2037
KEGGD05961 YesY
MeSHSulfur+dioxide
ChEBICHEBI:18422 YesY
ChEMBLCHEMBL1235997 N
RTECS numberWS4550000
Beilstein Reference3535237
Gmelin Reference1443
Jmol-3D imagesImage 1
Properties
Molecular formulaSO
2
Molar mass64.066 g mol−1
AppearanceColorless gas
Density2.6288 kg m−3
Melting point

-72 °C, 201 K, -98 °F

Boiling point

−10 °C, 263 K, 14 °F

Solubility in water94 g dm−3[1]
Vapor pressure237.2 kPa
Acidity (pKa)1.81
Basicity (pKb)12.19
Viscosity0.403 cP (at 0 °C)
Structure
Space groupC2v
Coordination
geometry
Digonal
Molecular shapeDihedral
Dipole moment1.62 D
Thermochemistry
Std enthalpy of
formation
ΔfHo298
-296.81 kJ mol−1
Standard molar
entropy
So298
248.223 J K−1 mol−1
Hazards
EU Index016-011-00-9
EU classificationToxic T
R-phrasesR23, R34, R50
S-phrases(S1/2), S9, S26, S36/37/39, S45
NFPA 704
NFPA 704.svg
0
3
0
LD503000 ppm (30 min inhaled, mouse)
Related compounds
Related sulfur oxidesSulfur monoxide
Sulfur trioxide
Related compoundsOzone

Selenium dioxide
Sulfurous acid
Tellurium dioxide

 N (verify) (what is: YesY/N?)
Except where noted otherwise, data are given for materials in their standard state (at 25 °C, 100 kPa)
Infobox references
 
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Sulfur dioxide
Identifiers
CAS number7446-09-5 YesY
PubChem1119
ChemSpider1087 YesY
UNII0UZA3422Q4 YesY
EC number231-195-2
UN number1079, 2037
KEGGD05961 YesY
MeSHSulfur+dioxide
ChEBICHEBI:18422 YesY
ChEMBLCHEMBL1235997 N
RTECS numberWS4550000
Beilstein Reference3535237
Gmelin Reference1443
Jmol-3D imagesImage 1
Properties
Molecular formulaSO
2
Molar mass64.066 g mol−1
AppearanceColorless gas
Density2.6288 kg m−3
Melting point

-72 °C, 201 K, -98 °F

Boiling point

−10 °C, 263 K, 14 °F

Solubility in water94 g dm−3[1]
Vapor pressure237.2 kPa
Acidity (pKa)1.81
Basicity (pKb)12.19
Viscosity0.403 cP (at 0 °C)
Structure
Space groupC2v
Coordination
geometry
Digonal
Molecular shapeDihedral
Dipole moment1.62 D
Thermochemistry
Std enthalpy of
formation
ΔfHo298
-296.81 kJ mol−1
Standard molar
entropy
So298
248.223 J K−1 mol−1
Hazards
EU Index016-011-00-9
EU classificationToxic T
R-phrasesR23, R34, R50
S-phrases(S1/2), S9, S26, S36/37/39, S45
NFPA 704
NFPA 704.svg
0
3
0
LD503000 ppm (30 min inhaled, mouse)
Related compounds
Related sulfur oxidesSulfur monoxide
Sulfur trioxide
Related compoundsOzone

Selenium dioxide
Sulfurous acid
Tellurium dioxide

 N (verify) (what is: YesY/N?)
Except where noted otherwise, data are given for materials in their standard state (at 25 °C, 100 kPa)
Infobox references

Sulfur dioxide (also sulphur dioxide) is the chemical compound with the formula SO
2
. At standard atmosphere it is a toxic gas with a pungent, irritating and rotten smell. The triple point is 197.69 K and 1.67Kpa. It is released naturally by volcanic activity and is a potent global warming gas.

In Earth's atmosphere, it exists in trace amounts of 1 part per billion by volume (ppbv), primarily as a pollutant commonly released by 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.

The specific person who discovered sulfur dioxide is not known, but it was first used by the Romans in winemaking, when they discovered that burning sulfur candles inside empty wine vessels keeps them fresh and free from vinegar smell.

Structure and bonding[edit]

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 +1.

The Lewis structure of sulfur dioxide consists of an S=O double bond and an S–O dative bond without utilizing d-orbitals,[3] resulting in a bond order of 1.5.

Occurrence[edit]

It is found on Earth and exists in very small concentrations and in the atmosphere at approximately 1 ppbv.[4][5]

On other planets, it can be found in various concentrations, the most significant being the atmosphere of Venus, it is the third most significant atmospheric gas, at 150 ppm and condenses to form clouds and is a key component of chemical reactions in the planet's atmosphere and contributes to global warming.[6] It has been implicated as a key agent in the warming of early Mars, with estimates of concentrations in the lower atmosphere as high as 100ppm[7] though it currently only exists in trace amounts. On both Venus and Mars, its primary source, like on Earth, is believed to be volcanic. It is also believed to exist in trace amounts in the atmosphere of Jupiter and atmosphere of Saturn.

As an ice, it is thought to exist in abundance on the Galilean moons – as sublimating ice or frost on the trailing hemisphere of Io, a natural satellite of Jupiter[8] and in the crust and mantle of Europa, Ganymede and Callisto possibly also in liquid form and readily reacting with water.[9]

Production[edit]

Sulfur dioxide is primarily produced for sulfuric acid manufacture (see contact process). In the United States in 1979, 23.6 million tonnes of sulfur dioxide was used in this way, compared with 150 thousand tonnes used for other purposes. Most sulfur dioxide is produced by the combustion of elemental sulfur. Some sulfur dioxide is also produced by roasting pyrite and other sulfide ores in air.[10]

Combustion routes[edit]

Sulfur dioxide is the product of the burning of sulfur or of burning materials that contain sulfur:

S + O2 → SO2, ΔH = -297 kJ/mol

To aid combustion, liquefied sulfur (140–150 °C) is sprayed through an atomizing nozzle to generate fine drops of sulfur with a large surface area. The reaction is exothermic, and the combustion produces temperatures of 1000–1600 °C. The significant amount of heat produced is recovered by steam generation that can subsequently be converted to electricity.[10]

The combustion of hydrogen sulfide and organosulfur compounds proceeds similarly. For example:

2 H2S + 3 O2 → 2 H2O + 2 SO2

The roasting of sulfide ores such as pyrite, sphalerite, and cinnabar (mercury sulfide) also releases SO2:[11]

4 FeS2 + 11 O2 → 2 Fe2O3 + 8 SO2
2 ZnS + 3 O2 → 2 ZnO + 2 SO2
HgS + O2 → Hg + SO2
4 FeS + 7O2 → 2 Fe2O3 + 4 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[edit]

Sulfur dioxide can also be 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

Up until the 1970s, commercial quantities of sulfuric acid and cement were produced by this process in Whitehaven Britain. Upon being mixed with shale or marl, and roasted, the sulfate liberated sulfur dioxide gas, used in sulfuric acid production, the reaction also produced calcium silicate, a precursor in cement production.[12]

On a laboratory scale, the action of hot sulfuric acid on copper turnings produces sulfur dioxide.

Cu + 2 H2SO4 → CuSO4 + SO2 + 2 H2O

From sulfite[edit]

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[edit]

Industrial reactions[edit]

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[edit]

Sulfur dioxide is one of the few common acidic yet reducing gases. It turns moist litmus pink (being acidic), then white (due to its bleaching effect). It may be identified by bubbling it through a dichromate solution, turning the solution from orange to green (Cr3+ (aq)). It can also reduce ferric ions to ferrous:

2Fe3+ + SO2 + 2H2O → 2Fe2+ + SO42− +4H+

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.[13]

Uses[edit]

Precursor to sulfuric acid[edit]

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[edit]

Sulfur dioxide is sometimes used as a preservative for dried apricots, dried figs, 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 colorful appearance of the fruit and prevents rotting. It is also added to sulfured molasses.

In winemaking[edit]

The specific person who discovered sulfur dioxide is not known, but it was first used by the Romans in winemaking, when they discovered that burning sulfur candles inside empty wine vessels keeps them fresh and free from vinegar smell.[14]

Sulfur dioxide is an important compound in winemaking, and is designated as parts per million in wine, E number: E220.[15] It is present even in so-called unsulfurated wine at concentrations of up to 10 mg/L.[16] It serves as an antibiotic and antioxidant, protecting wine from spoilage by bacteria and oxidation. Its antimicrobial action also helps to minimize 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 combinations 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 pH more SO2 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 50 ppm, 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 due the risk of cork taint,[17] 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[edit]

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.[18]

Sulfur dioxide is fairly soluble in water, and by both IR and Raman spectroscopy, it is known that the hypothetical sulfurous acid, H2SO3, is not present to any extent. However, such solutions do show spectra of the hydrogen sulfite ion, HSO3, by reaction with water, and it is in fact the actual reducing agent present:

SO2 + H2O HSO3 + H+

Biochemical and biomedical roles[edit]

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.[19] Sulfur dioxide blocks nerve signals from the pulmonary stretch receptors (PSRs) and abolishes the Hering–Breuer inflation reflex.

As a refrigerant[edit]

Being easily condensed and possessing a high heat of evaporation, sulfur dioxide is a candidate material for refrigerants. Prior to the development of CFCs, sulfur dioxide was used as a refrigerant in home refrigerators.

As a reagent and solvent in the laboratory[edit]

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:[20]

Preparation of m-trifluoromethylbenzenesulfonyl chloride.svg

As an air pollutant[edit]

A sulfur dioxide plume from the Halemaʻumaʻu vent, glows at night

Sulfur dioxide is a noticeable component in the atmosphere, especially following volcanic eruptions.[21] According to the United States Environmental Protection Agency (as presented by the 2002 World Almanac or in chart form[22]), the following amount of sulfur dioxide was released in the U.S. per year:

YearSO2
197031,161,000 short tons (28.3 Mt)
198025,905,000 short tons (23.5 Mt)
199023,678,000 short tons (21.5 Mt)
199618,859,000 short tons (17.1 Mt)
199719,363,000 short tons (17.6 Mt)
199819,491,000 short tons (17.7 Mt)
199918,867,000 short tons (17.1 Mt)

Sulfur dioxide is a major air pollutant and has significant impacts upon human health.[23] In addition the concentration of sulfur dioxide in the atmosphere can influence the habitat suitability for plant communities as well as animal life.[24] 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% 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.[25]

Sulfur can also be removed from fuels prior to burning the fuel, preventing 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. Redox processes using iron oxides can also be used, for example, Lo-Cat[26] or Sulferox.[27]

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.[28]

As of 2006, China was the world's largest sulfur dioxide polluter, with 2005 emissions estimated to be 25,490,000 short tons (23.1 Mt). This amount represents a 27% increase since 2000, and is roughly comparable with U.S. emissions in 1980.[29]

Safety[edit]

Inhalation[edit]

Inhaling sulfur dioxide is associated with increased respiratory symptoms and disease, difficulty in breathing, and premature death.[30] 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.[31]

A 2011 systematic review concluded that exposure to sulfur dioxide is associated with preterm birth.[32]

Ingestion[edit]

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 asthmatic individuals who may be sensitive to them, especially in large amounts.[33] Symptoms of sensitivity to sulfiting agents, including sulfur dioxide, manifest as potentially life-threatening trouble breathing within minutes of ingestion.[34]

See also[edit]

References[edit]

  1. ^ Lide, David R., ed. (2006). CRC Handbook of Chemistry and Physics (87th ed.). Boca Raton, FL: CRC Press. ISBN 0-8493-0487-3. 
  2. ^ Holleman, A. F.; Wiberg, E. (2001), Inorganic Chemistry, San Diego: Academic Press, ISBN 0-12-352651-5 
  3. ^ Cunningham, Terence P.; Cooper, David L.; Gerratt, Joseph; Karadakov, Peter B. and Raimondi, Mario (1997). "Chemical bonding in oxofluorides of hypercoordinatesulfur". Journal of the Chemical Society, Faraday Transactions 93 (13): 2247–2254. doi:10.1039/A700708F. 
  4. ^ Owen, Lewis A.; Pickering, Kevin T (1997). An Introduction to Global Environmental Issues. Taylor & Francis. pp. 33–. ISBN 978-0-203-97400-1. 
  5. ^ Taylor, J.A.; Simpson, R.W.; Jakeman, A.J. (1987). "A hybrid model for predicting the distribution of sulphur dioxide concentrations observed near elevated point sources". Ecological Modelling 36 (3–4): 269–296. doi:10.1016/0304-3800(87)90071-8. ISSN 0304-3800. 
  6. ^ Marcq, Emmanuel; Bertaux, Jean-Loup; Montmessin, Franck; Belyaev, Denis (2012). "Variations of sulphur dioxide at the cloud top of Venus’s dynamic atmosphere". Nature Geoscience. doi:10.1038/ngeo1650. ISSN 1752-0894. 
  7. ^ Halevy, I.; Zuber, M. T.; Schrag, D. P. (2007). "A Sulfur Dioxide Climate Feedback on Early Mars". Science 318 (5858): 1903–1907. doi:10.1126/science.1147039. ISSN 0036-8075. 
  8. ^ Cruikshank, D. P.; Howell, R. R.; Geballe, T. R.; Fanale, F. P. (1985). Sulfur Dioxide Ice on IO. pp. 805–815. doi:10.1007/978-94-009-5418-2_55. 
  9. ^ Europa's Hidden Ice Chemistry – NASA Jet Propulsion Laboratory. Jpl.nasa.gov (2010-10-04). Retrieved on 2013-09-24.
  10. ^ a b Müller, Hermann (2005), "Sulfur Dioxide", Ullmann's Encyclopedia of Industrial Chemistry, Weinheim: Wiley-VCH, doi:10.1002/14356007.a25_569 
  11. ^ Shriver, Atkins. Inorganic Chemistry, Fifth Edition. W. H. Freeman and Company; New York, 2010; p. 414.
  12. ^ WHITEHAVEN COAST ARCHAEOLOGICAL SURVEY. lakestay.co.uk (2007)
  13. ^ Greenwood, Norman N.; Earnshaw, Alan (1997). Chemistry of the Elements (2nd ed.). Butterworth-Heinemann. ISBN 0080379419. 
  14. ^ "Practical Winery & vineyard Journal Jan/Feb 2009". www.practicalwinery.com. 1 Feb 2009. 
  15. ^ Current EU approved additives and their E Numbers, The Food Standards Agency website.
  16. ^ Sulphites in wine, MoreThanOrganic.com.
  17. ^ Chlorine Use in the Winery. Purdue University
  18. ^ Tchobanoglous, George (1979). Wastewater Engineering (3rd ed.). New York: McGraw Hill. ISBN 0-07-041677-X. 
  19. ^ Liu, D.; Jin, H; Tang, C; Du, J (2010). "Sulfur dioxide: a novel gaseous signal in the regulation of cardiovascular functions". Mini-Reviews in Medicinal Chemistry 10 (11): 1039–1045. PMID 20540708. 
  20. ^ Hoffman, R. V. (1990), "m-Trifluoromethylbenzenesulfonyl Chloride", Org. Synth. ; Coll. Vol. 7: 508 
  21. ^ Volcanic Gases and Their Effects. Volcanoes.usgs.gov. Retrieved on 2011-10-31.
  22. ^ National Trends in Sulfur Dioxide Levels, United States Environmental Protection Agency.
  23. ^ Sulfur Dioxide. United States Environmental Protection Agency
  24. ^ Hogan, C. Michael (2010). "Abiotic factor" in Encyclopedia of Earth. Emily Monosson and C. Cleveland (eds.). National Council for Science and the Environment. Washington DC
  25. ^ 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. 
  26. ^ FAQ’s About Sulfur Removal and Recovery using the LO-CAT® Hydrogen Sulfide Removal System. gtp-merichem.com
  27. ^ Process screening analysis of alternative gas treating and sulfur removal for gasification. (December 2002) Report by SFA Pacific, Inc. prepared for U.S. Department of Energy (PDF) . Retrieved on 2011-10-31.
  28. ^ May, Walter R. Marine Emissions Abatement. SFA International, Inc., p. 6.
  29. ^ China has its worst spell of acid rain, United Press International (2006-09-22).
  30. ^ Sulfur Dioxide U.S. Environmental Protection Agency
  31. ^ "NIOSH Pocket Guide to Chemical Hazards". 
  32. ^ Shah PS, Balkhair T, Knowledge Synthesis Group on Determinants of Preterm/LBW Births (2011). "Air pollution and birth outcomes: a systematic review". Environ Int 37 (2): 498–516. doi:10.1016/j.envint.2010.10.009. PMID 21112090. 
  33. ^ "Center for Science in the Public Interest – Chemical Cuisine". Retrieved March 17, 2010. 
  34. ^ "California Department of Public Health: Food and Drug Branch: Sulfites". Retrieved September 27, 2013. 

External links[edit]