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Sulfur dioxide

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Sulfur dioxide
Skeletal formula sulfur dioxide with assorted dimensions
Spacefill model of sulfur dioxide
Names
IUPAC name
Sulfur dioxide
Other names
Sulfurous anhydride
Sulfur(IV) oxide
Identifiers
3D model (JSmol)
3535237
ChEBI
ChEMBL
ChemSpider
ECHA InfoCard 100.028.359 Edit this at Wikidata
EC Number
  • 231-195-2
E number E220 (preservatives)
1443
KEGG
MeSH Sulfur+dioxide
RTECS number
  • WS4550000
UNII
UN number 1079, 2037
  • InChI=1S/O2S/c1-3-2 checkY
    Key: RAHZWNYVWXNFOC-UHFFFAOYSA-N checkY
  • InChI=1/O2S/c1-3-2
    Key: RAHZWNYVWXNFOC-UHFFFAOYAT
  • O=S=O
Properties
SO
2
Molar mass 64.066 g mol−1
Appearance Colorless gas
Odor Just-struck match/eggs like
Density 2.6288 kg m−3
Melting point −72 °C; −98 °F; 201 K
Boiling point −10 °C (14 °F; 263 K)
94 g/L[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
248.223 J K−1 mol−1
-296.81 kJ mol−1
Hazards
NFPA 704 (fire diamond)
NFPA 704 four-colored diamondHealth 3: Short exposure could cause serious temporary or residual injury. E.g. chlorine gasFlammability 0: Will not burn. E.g. waterInstability 0: Normally stable, even under fire exposure conditions, and is not reactive with water. E.g. liquid nitrogenSpecial hazards (white): no code
3
0
0
Lethal dose or concentration (LD, LC):
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).
☒N verify (what is checkY☒N ?)

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.

Sulfur dioxide was used by the Romans in winemaking, when they discovered that burning sulfur candles inside empty wine vessels kept them fresh and free from vinegar smell.[2]

Structure and bonding

SO2 is a bent molecule with C2v symmetry point group. A valence bond theory approach which just considered s and p orbitals would describe the bonding in terms of resonance between two resonance structures.

Two resonance structures of sulfur dioxide

The sulfur - oxygen bond has a bond order of 1.5. There is support for this simple approach that does not invoke d orbital participation.[3] In terms of electron-counting formalism, the sulfur atom has an oxidation state of +4 and a formal charge of +1.

Occurrence

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

On other planets, it can be found in various concentrations, the most significant being the atmosphere of Venus, where it is the third-most significant atmospheric gas at 150 ppm. There, it 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 100 ppm,[7] though it 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.

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

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 were 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

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

Sulfur dioxide can also be a byproduct 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

Until the 1970s, commercial quantities of sulfuric acid and cement were produced by this process in Whitehaven, England. 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

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 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 [citation needed]

Sulfur dioxide can react with certain 1,3-dienes in a cheletropic reaction to form cyclic sulfones. This reaction is exploited on an industrial scale for the synthesis of sulfolane, which is an important solvent in the petrochemical industry.

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

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, dried figs, and other dried fruits, owing to its antimicrobial properties, and is called E220.[14] when used in this way in Europe. As a preservative, it maintains the colorful appearance of the fruit and prevents rotting. It is also added to sulfured molasses.

In winemaking

Sulfur dioxide was 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.[15]

Sulfur dioxide is still an important compound in winemaking, and is measured in parts per million in wine 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. The molecular SO2 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

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; 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

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 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 chlorofluorocarbons, 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 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 yields the corresponding aryl sulfonyl chloride, for example:[20]

As an air pollutant

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,[22] the amount of sulfur dioxide released in the U.S. per year was:

Year SO2
1970 31,161,000 short tons (28.3 Mt)
1980 25,905,000 short tons (23.5 Mt)
1990 23,678,000 short tons (21.5 Mt)
1996 18,859,000 short tons (17.1 Mt)
1997 19,363,000 short tons (17.6 Mt)
1998 19,491,000 short tons (17.7 Mt)
1999 18,867,000 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 had 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 no sulfur remains 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 used in gasoline and diesel engines 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

Inhalation

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 to 5 parts per million (ppm). The OSHA PEL is currently set at 2 ppm (5.64  mg/m3) time-weighted average. NIOSH has set the IDLH at 100 ppm.[31] In 2010, the EPA "revised the primary SO2 NAAQS by establishing a new one-hour standard at a level of 75 parts per billion (ppb). EPA revoked the two existing primary standards because they would not provide additional public health protection given a one-hour standard at 75 ppb."[32]

A 2011 systematic review concluded that exposure to sulfur dioxide is associated with preterm birth.[33]

Ingestion

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.[34] Symptoms of sensitivity to sulfiting agents, including sulfur dioxide, manifest as potentially life-threatening trouble breathing within minutes of ingestion.[35]

See also

References

  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. ^ "Practical Winery & Vineyard Journal Jan/Feb 2009". www.practicalwinery.com. 1 Feb 2009.
  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.{{cite journal}}: CS1 maint: multiple names: authors list (link)
  4. ^ Owen, Lewis A.; Pickering, Kevin T (1997). An Introduction to Global Environmental Issues. Taylor & Francis. pp. 33–. ISBN 978-0-203-97400-1.{{cite book}}: CS1 maint: multiple names: authors list (link)
  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": 805–815. doi:10.1007/978-94-009-5418-2_55. {{cite journal}}: Cite journal requires |journal= (help)
  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. "Sulfur Dioxide". Ullmann's Encyclopedia of Industrial Chemistry. Weinheim: Wiley-VCH. doi:10.1002/14356007.a25_569. ISBN 978-3527306732.
  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 978-0-08-037941-8.
  14. ^ Current EU approved additives and their E Numbers, The Food Standards Agency website.
  15. ^ "Practical Winery & vineyard Journal Jan/Feb 2009". www.practicalwinery.com. 1 Feb 2009.
  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". Organic Syntheses; Collected Volumes, vol. 7, p. 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. ^ http://www.epa.gov/airquality/sulfurdioxide/
  33. ^ 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.{{cite journal}}: CS1 maint: multiple names: authors list (link)
  34. ^ "Center for Science in the Public Interest – Chemical Cuisine". Retrieved March 17, 2010.
  35. ^ "California Department of Public Health: Food and Drug Branch: Sulfites" (PDF). Retrieved September 27, 2013.
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