Sulfur dioxide
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2008-07-06T15:23:05Z
TGC55
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/* Emissions */ link to Halemaumau Crater
{{redirect|SO2|the video game|Star Ocean: The Second Story}}
{{Chembox new
| Name = Sulfur dioxide
| ImageFileL1 = Sulfur-dioxide-2D.svg
| ImageSizeL1 = 120
| ImageFileR1 = Sulfur-dioxide-3D-vdW.png
| OtherNames = sulfur(IV) oxide; sulfurous anhydride
| Section1 = {{Chembox Identifiers
| CASNo = 7446-09-5
| RTECS = WS4550000
}}
| Section2 = {{Chembox Properties
| Formula = [[Sulfur|S]][[Oxygen|O]]<sub>2</sub>
| MolarMass = 64.054 g mol<sup>−1</sup>
| Appearance = colourless gas
| Density = 2.551 g/L, gas
| Solubility = 9.4 g/100 mL (25 °C)
| MeltingPt = −72.4 °C (200.75 K)
| BoilingPt = −10 °C (263 K)
| pKa = 1.81
}}
| Section3 = {{Chembox Structure
| MolShape = Bent 120°<ref>[http://chemed.chem.purdue.edu/genchem/topicreview/bp/ch8/table.php Table of Geometries based on VSEPR<!-- Bot generated title -->]</ref>
| Dipole = 1.63 [[Debye|D]]
}}
| Section7 = {{Chembox Hazards
| FlashPt = non-flammable
| EUClass = Toxic
| RPhrases = {{R23}} {{R34}}
| SPhrases = {{S1/2}} {{S9}} {{S26}} {{S36/37/39}} {{S45}}
| NFPA-H = 3
}}
| Section8 = {{Chembox Related
| OtherCpds = [[Sulfur trioxide]]; [[sulfuric acid]]
}}
}}
'''Sulfur dioxide''' (also '''sulphur dioxide''') is the [[chemical compound]] with the formula SO<sub>2</sub>. SO<sub>2</sub> is produced by [[volcano]]es and in various industrial processes. Since [[coal]] and [[petroleum]] often contain sulfur compounds, their combustion generates sulfur dioxide. Further oxidation of SO<sub>2</sub>, usually in the presence of a catalyst such as [[nitrogen dioxide|NO<sub>2</sub>]], forms [[sulfuric acid|H<sub>2</sub>SO<sub>4</sub>]], and thus [[acid rain]].<ref>Holleman, A. F.; Wiberg, E. "Inorganic Chemistry" Academic Press: San Diego, 2001. ISBN 0-12-352651-5.</ref> This is one of the causes for concern over the environmental impact of the use of these fuels as power sources.
==Preparation==
Sulfur dioxide can be prepared by burning [[sulfur]]:
:S<sub>8</sub> + 8 O<sub>2</sub> → 8 SO<sub>2</sub>
The combustion of [[hydrogen sulfide]] and organosulfur compounds proceeds similarly.
:2 H<sub>2</sub>S(g) + 3 O<sub>2</sub>(g) → 2 H<sub>2</sub>O(g) + 2 SO<sub>2</sub>(g)
The roasting of sulfide ores such as iron [[pyrites]], [[sphalerite]] (zinc blende) and [[cinnabar]] (mercury sulfide) also releases SO<sub>2</sub>:
:4 [[Iron|Fe]]S<sub>2</sub>(s) + 11 O<sub>2</sub>(g) → 2 Fe<sub>2</sub>O<sub>3</sub>(s) + 8 SO<sub>2</sub>(g)
:2 [[Zinc|Zn]]S(s) + 3 O<sub>2</sub>(g) → 2 ZnO(s) + 2 SO<sub>2</sub>(g)
:HgS(s) + O<sub>2</sub>(g) → Hg(g) + SO<sub>2</sub>(g)
Sulfur dioxide is a by-product in the manufacture of [[cement]]: [[Calcium silicate|CaSiO<sub>3</sub>]] and [[Calcium sulfate|CaSO<sub>4</sub>]] is heated with [[coke (fuel)|coke]] and sand in this process:
:2 CaSO<sub>4</sub>(s) + 2SiO<sub>2</sub>(s) + C(s) → 2 CaSiO<sub>3</sub>(s) + 2 SO<sub>2</sub>(g) + CO<sub>2</sub>(g)
Action of hot [[sulfuric acid]] on copper [[swarf|turnings]] produces sulfur dioxide.
:Cu(s) + 2H<sub>2</sub>SO<sub>4</sub>(aq) → CuSO<sub>4</sub>(aq) + SO<sub>2</sub>(g) + 2H<sub>2</sub>O(l)
<<<<
==Structure and bonding==
SO<sub>2</sub> is a bent molecule with C<sub>2v</sub> [[symmetry point group]]. In terms of [[electron counting|electron-counting]] formalisms, the sulfur atom has an [[oxidation state]] of +4, a [[formal charge]] of 0, and is surrounded by 5 [[electron pair]]s 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.
<center>
{|align="center" class="wikitable"
|<center>[[Image:Sulfur-dioxide-resonance-2D.png|500px]]</center>
|-
|<center>three [[resonance (chemistry)|resonance structures]] of sulfur dioxide</center>
|}
</center>
The S-O bonds are shorter in SO<sub>2</sub> (143.1 pm) than in [[sulfur monoxide]], SO (148.1 pm), whereas the O-O bonds are longer in O<sub>3</sub> (127.8 pm) than in [[dioxygen]], O<sub>2</sub> (120.7 pm). The mean bond energy is greater in SO<sub>2</sub> (548 kJ mol<sup>−1</sup>) than in SO (524 kJ mol<sup>−1</sup>), whereas its is less in O<sub>3</sub> (297 kJ mol<sup>−1</sup>) than in O<sub>2</sub> (490 kJ mol<sup>−1</sup>). 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<ref>{{Greenwood&Earnshaw}} p. 700</ref>
==Reactions==
Treatment of basic solutions with sulfur dioxide affords [[sulfite]] salts:
:SO<sub>2</sub> + 2 NaOH → Na<sub>2</sub>SO<sub>3</sub> + H<sub>2</sub>O
Featuring sulfur in the +4 oxidation state, sulfur dioxide is a [[reducing agent]]. It is oxidized by halogens such as chlorine to give the sulfuryl halides:
:SO<sub>2</sub> + Cl<sub>2</sub> → [[Sulfuryl chloride|SO<sub>2</sub>Cl<sub>2</sub>]]
However, on rare occasions, it can also act as an [[oxidising agent]]: in the [[Claus process]], sulfur dioxide is reduced by hydrogen sulfide to give elemental sulfur:
:SO<sub>2</sub> + 2 H<sub>2</sub>S → 3 S + 2 H<sub>2</sub>O
Sulfur dioxide can act as a metal binding [[ligand]], typically where the transition metal is in oxidation state 0 or +1.<ref name = "Greenwood">{{Greenwood&Earnshaw}}</ref> Up to 9 different bonding modes have been determined which include<ref name = "Greenwood"/>:
*S donation -planar and pyramidal [[hapticity|η]]<sup>1</sup>
*O donation [[hapticity|η]]<sup>1</sup>
*π donation- side on [[hapticity|η]]<sup>2</sup>
*S bridging across two metal centres or two ends of a metal-metal bon
*O-S bridging to two metal centres
*bridging- one metal centre π donation- side on, the other metal center O donation
*bridging over three metal centres
==Uses==
===As a preservative===
Sulfur dioxide is sometimes used as a [[preservative]] for dried [[apricot]]s and other [[dried fruit]]s due to its [[antimicrobial]] properties, it is sometimes called [[E number|E220]] when used in this way. The preservative is used to maintain the appearance of the fruit and prevent [[Decomposition|rotting]]. Its presence can give fruit a distinctive chemical [[taste]].
===In winemaking===
Sulfur dioxide is a very important compound in winemaking, and is designated as parts per million in wine, [[E number]]: E220.<ref>[http://www.food.gov.uk/safereating/chemsafe/additivesbranch/enumberlist#h_3 Current EU approved additives and their E Numbers], The Food Standards Agency website.</ref> It is present even in so-called unsulphurated wine at concentrations of up to 10 milligrams per litre.<ref>[http://www.morethanorganic.com/sulphur-in-the-bottle Sulphites in wine], MoreThanOrganic.com.</ref> It serves as antioxidant, protecting wine from spoilage by bacteria and oxidation. It also helps to keep volatile acidity at desirable levels. Sulfur dioxide is responsible for the words "contains sulfites" found on wine labels. Wines with SO<sub>2</sub> concentrations below 10ppm do not require "contains sulfites" on the label by US and EU laws. The upper limit of SO<sub>2</sub> allowed in wine is 350ppm in US, in the EU is 160 ppm for red wines and 210 ppm for white and [[rosé|rosé wines]]. In low concentrations SO<sub>2</sub> is mostly undetectable in wine, but at over 50ppm, SO<sub>2</sub> becomes evident in the nose and taste of wine.
SO<sub>2</sub> is also a very important element in winery sanitation. Wineries and equipment must be kept very clean, and because bleach cannot be used in a winery, a mixture of SO<sub>2</sub>, water, and citric acid is commonly used to clean hoses, tanks, and other equipment to keep it clean and free of bacteria.
===As a reducing bleach===
Sulfur dioxide is also a good [[Reducing agent|reductant]]. In the presence of water, sulfur dioxide is able to decolorize substances. Specifically it is a useful reducing [[bleach]] for [[paper]]s 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.
===Precursor to sulfuric acid===
Sulfur dioxide is also used to make 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]].
===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. Sulfur dioxide has no role in mammalian biology. Sulfur dioxide blocks nerve signals from the [[pulmonary stretch receptors]] (PSR's) and abolishes the [[Hering-Breuer reflex|Hering-Breuer inflation reflex]].
===As a refrigerant===
Being easily condensed and with a high heat of evaporation, sulfur dioxide is a candidate material for refrigerants. Prior to the development of [[Haloalkane|freons]], sulfur dioxide was used as a [[refrigerant]] in home refrigerators.
===As a reagent and solvent===
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 salt]]s with sulfur dioxide affords the corresponding aryl sulfonyl chloride.<ref>R. V. Hoffman “m-Trifluoromethylbenzenesulfonyl chloride” Organic Syntheses, Collected Volume 7, p.508 (1990). http://www.orgsyn.org/orgsyn/pdfs/CV7P0508.pdf.</ref>
===Dechlorination===
In municipal wastewater treatment sulfur dioxide is used to treat chlorinated wastewater prior to release. Sulfur dioxide reacts with free and combined chlorine to form negatively charged chloride ions. <ref>Tchobanoglous, George. Wastewater Engineering. 3rd ed. New York: Mc Graw Hill, 1979.</ref>
==Emissions==
[[Image:Sulfur dioxide emissions from the Halema`uma`u vent, glows at night.jpg|Sulfur dioxide emissions from the [[Halemaumau Crater|Halema`uma`u vent]], glows at night|thumb|300px]]
According to the [[United States Environmental Protection Agency|U.S. EPA]] (as presented by the ''2002 World Almanac'' or in chart form<ref>[http://www.epa.gov/air/airtrends/sulfur.html National Trends in Sulfur Dioxide Levels], [[United States Environmental Protection Agency]].</ref>), the following amount of sulfur dioxide was released in the [[U.S.]] per year, measured in thousands of [[short tons]]:
{| border=1
|*'''1999'''
|18,867
|-
|*'''1998'''
|19,491
|-
|*'''1997'''
|19,363
|-
|*'''1996'''
|18,859
|-
|*'''1990'''
|23,678
|-
|*'''1980'''
|25,905
|-
|*'''1970'''
|31,161
|}
Due largely to the [[United States|US]] EPA’s [[Acid Rain Program]], the U.S. has witnessed a 33 percent decrease in emissions between 1983 and 2002. This improvement resulted from [[flue gas desulfurization]], a technology that enables SO<sub>2</sub> to be chemically bound in [[power plant]]s burning sulfur-containing [[coal]] or [[Petroleum|oil]]. In particular, [[calcium oxide|calcium oxide (lime)]] reacts with sulfur dioxide to form [[calcium sulfite]]:
:CaO + SO<sub>2</sub> → CaSO<sub>3</sub>
Aerobic oxidation<!--right?--> converts this CaSO<sub>3</sub> into CaSO<sub>4</sub>, [[gypsum]]. Most gypsum sold in Europe comes from flue gas desulfurization.
New fuel additive catalysts, such as [[ferox (fuel additive)|ferox]], are being used in gasoline and diesel engines in order to lower the emission of sulfur oxide gases into the atmosphere. This is also done by forcing the sulfur into stable mineral salts and mixed mineral sulfates as opposed to sulfuric acid and sulfur oxides.
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<ref>[http://upi.com/NewsTrack/view.php?StoryID=20060922-010840-7049r China has its worst spell of acid rain], [[United Press International]].</ref>.
[[Al-Mishraq]], an Iraqi sulfur plant, was the site of a 2003 disaster resulting in the release of massive amounts of sulfur dioxide into the atmosphere.
==Temperature dependence of aqueous solubility==
{|border=1
|22 g/100ml (0 °C)
|15 g/100ml (10 °C)
|-
|11 g/100ml (20 °C)
|9.4 g/100 ml (25 °C)
|-
|8 g/100ml (30 °C)
|6.5 g/100ml (40 °C)
|-
|5 g/100ml (50 °C)
|4 g/100ml (60 °C)
|-
|3.5 g/100ml (70 °C)
|3.4 g/100ml (80 °C)
|-
|3.5 g/100ml (90 °C)
|3.7 g/100ml (100 °C)
|}
* The values are tabulated for 101.3 kPa [[partial pressure]] of SO<sub>2</sub>. [[Solubility]] of gas in a liquid depends on the gas [[partial pressure]] according to [[Henry's law]].
* The solublity is given for "pure water", i.e., water that contains only SO<sub>2</sub> in the amount at equilibrium with the gas phase. This "pure water" is going to be acidic. The solublity of SO<sub>2</sub> in neutral (or alkaline) water is generally going to be higher because of the [[pH]]-dependent speciation of SO<sub>2</sub> in the solution with the production of [[bisulfite]] and some [[sulfite]] ions.
==Threats to Health==
Sulfur dioxide acts as an acid. Inhalation results in labored breathing, coughing, and/or a sore throat and may cause permanent pulmonary damage. When mixed with water and contacted by skin, frostbite may occur. When it makes contact with eyes, redness and pain will occur.<ref> [http://www.ilo.org/public/english/protection/safework/cis/products/icsc/dtasht/_icsc00/icsc0074.htm SULPHUR DIOXIDE], [[International Labour Organization]].</ref>
==References==
{{Reflist}}
==See also==
* [[Sulfur trioxide]]
* [[Sulfur-iodine cycle]]
* [[National Ambient Air Quality Standards]]
* [[Homer City Generating Station]]
*
{{commonscat|sulfur dioxide}}
<123>
==External links==
*[http://www.epa.gov/air/airtrends/sulfur.html United States Environmental Protection Agency Sulfur Dioxide page]
*[http://www.ilo.org/public/english/protection/safework/cis/products/icsc/dtasht/_icsc00/icsc0074.htm International Chemical Safety Card 0074]
*[http://www-cie.iarc.fr/htdocs/monographs/vol54/02-sulfur-dioxide.htm IARC Monograph "Sulfur Dioxide and some Sulfites, Bisulfites and Metabisulfites"]
*[http://www.cdc.gov/niosh/npg/npgd0575.html NIOSH Pocket Guide to Chemical Hazards]
*[http://www.fedupwithfoodadditives.info/factsheets/Factsulphites.htm Food Intolerance Network] - Sulfite factsheet
*[http://www.chm.bris.ac.uk/motm/so2/so2h.htm Sulfur Dioxide, Molecule of the Month]
[[Category:Sulfur oxides]]
[[Category:Acidic oxides]]
[[Category:Hazardous air pollutants]]
[[Category:Smog]]
[[Category:Preservatives]]
[[Category:Refrigerants]]
[[Category:IARC Group 3 carcinogens]]
[[als:Schwefeldioxid]]
[[bs:Sumpor dioksid]]
[[bg:Серен диоксид]]
[[ca:Diòxid de sofre]]
[[cs:Oxid siřičitý]]
[[cy:Sylffwr deuocsid]]
[[de:Schwefeldioxid]]
[[el:Διοξείδιο του θείου]]
[[es:Óxido de azufre (IV)]]
[[eo:Sulfura dioksido]]
[[fr:Dioxyde de soufre]]
[[hi:सल्फर डाइआक्साइड]]
[[hr:Sumporov dioksid]]
[[it:Diossido di zolfo]]
[[he:גופרית דו-חמצנית]]
[[lv:Sēra dioksīds]]
[[lt:Sieros dioksidas]]
[[hu:Kén-dioxid]]
[[mk:Сулфур диоксид]]
[[nl:Zwaveldioxide]]
[[ja:二酸化硫黄]]
[[no:Svoveldioksid]]
[[pl:Tlenek siarki(IV)]]
[[pt:Dióxido de enxofre]]
[[ru:Оксид серы(IV)]]
[[sk:Oxid siričitý]]
[[sl:Žveplov dioksid]]
[[sr:Сумпор диоксид]]
[[su:Walirang dioksida]]
[[fi:Rikkidioksidi]]
[[sv:Svaveldioxid]]
[[th:ซัลเฟอร์ไดออกไซด์]]
[[vi:Điôxít lưu huỳnh]]
[[uk:Діоксид сірки]]
[[zh:二氧化硫]]