Difference between revisions of "Sulfur dioxide"

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|  CASNo = 7446-09-5
 
|  CASNo = 7446-09-5
 
|    CASNo_Ref = {{cascite}}
 
|    CASNo_Ref = {{cascite}}
 +
|  EC-number = 231-195-2
 
|  ChemSpiderID = 1087
 
|  ChemSpiderID = 1087
 
|  RTECS = WS4550000
 
|  RTECS = WS4550000
−
|  EINECS = 231-195-2
 
 
   }}
 
   }}
 
| Section2 = {{Chembox Properties
 
| Section2 = {{Chembox Properties
 +
|  Reference = <ref>{{citation | first = Pradyot | last = Patnaik | title = Handbook of Inorganic Chemicals | publisher = McGraw-Hill | year = 2002 | isbn = 0070494398}}.</ref>
 
|  Formula = SO<sub>2</sub>
 
|  Formula = SO<sub>2</sub>
−
|  MolarMass = 64.07&nbsp;g/mol
+
|  MolarMass = 64.064 g/mol
 
|  Appearance = colorless gas
 
|  Appearance = colorless gas
−
|  Density = 2.551&nbsp;g/L (gas)<br/>1.46&nbsp;g/cm<sup>3</sup> {{nowrap|(liquid, −10 °C)}}
+
|  Density = 2.551 g/L (gas)<br/>1.46 g/cm<sup>3</sup> {{nowrap|(liquid, −10 °C)}}
−
|  Solubility = 22.97&nbsp;g/100 mL (0&nbsp;°C) <br> 11.58&nbsp;g/100&nbsp;mL (20&nbsp;°C) <br> 9.4&nbsp;g/100 mL (25&nbsp;°C) <ref>Pradyot Patnaik. ''Handbook of Inorganic Chemicals''. McGraw-Hill, 2002, ISBN 0070494398</ref>
+
|  Solubility = 9.4 g/100 mL (25 °C)
−
|  SolubleOther = very soluble in [[acetone]], [[methyl isobutyl ketone]], [[acetic acid]], [[alcohol]] <br> soluble in [[sulfuric acid]]
+
|  SolubleOther = very soluble in [[acetone]], [[methyl isobutyl ketone]], [[acetic acid]], [[alcohol]] <br/> soluble in [[sulfuric acid]]
 
|  MeltingPtC = -75.5
 
|  MeltingPtC = -75.5
 
|  BoilingPtC = -10.0
 
|  BoilingPtC = -10.0
 
|  pKa = 1.81
 
|  pKa = 1.81
−
|  Viscosity = 0.403&nbsp;[[Poise|cP]] (0&nbsp;°C)
+
|  Viscosity = 0.403 cP (0 °C)
 
   }}
 
   }}
 
| Section3 = {{Chembox Structure
 
| Section3 = {{Chembox Structure
Line 27: Line 28:
 
   }}
 
   }}
 
| Section7 = {{Chembox Hazards
 
| Section7 = {{Chembox Hazards
−
|  ExternalMSDS = [http://www.ilo.org/public/english/protection/safework/cis/products/icsc/dtasht/_icsc00/icsc0074.htm ICSC 0074]
+
|  ExternalMSDS = {{ICSC-small|00|74}}
 
|  EUIndex = 016-011-00-9
 
|  EUIndex = 016-011-00-9
−
|  EUClass = Toxic ('''T''')<br/>Corrosive ('''C''')
+
|  GHSPictograms =  
−
|  RPhrases = {{R23}} {{R34}}
+
|  GHSSignalWord =  
−
|  SPhrases = {{S1/2}} {{S9}} {{S26}} {{S36/37/39}} {{S45}}
+
|  HPhrases =  
−
|  NFPA-H = 3
+
|  FlashPt = non-flammable
−
|  NFPA-F = 0
+
|  LD50 = 3000 ppm (30&nbsp;min inhaled, mouse)
−
|  NFPA-R = 0
 
−
|  NFPA-O =
 
−
|  FlashPt = Non-flammable
 
−
|  LD50 = 3000&nbsp;ppm (30&nbsp;min inhaled, mouse)
 
 
|  PEL =
 
|  PEL =
 
   }}
 
   }}
 
| Section8 = {{Chembox Related
 
| Section8 = {{Chembox Related
 
|  OtherAnions =
 
|  OtherAnions =
−
|  OtherCations = [[Selenium dioxide]]<br/>[[Tellurium dioxide]]
+
|  OtherCations = [[Ozone]]<br/>[[Selenium dioxide]]<br/>[[Tellurium dioxide]]
 
|  OtherFunctn = [[Sulfur monoxide]]<br/>[[Sulfur trioxide]]
 
|  OtherFunctn = [[Sulfur monoxide]]<br/>[[Sulfur trioxide]]
 
|    Function = [[sulfur]] [[oxide]]s
 
|    Function = [[sulfur]] [[oxide]]s
Line 49: Line 46:
 
}}
 
}}
  
−
'''Sulfur dioxide''' (also '''sulphur dioxide''') is the [[chemical compound]] with the formula SO<sub>2</sub>.  It 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&Wiberg}}</ref> This is one of the causes for concern over the environmental impact of the use of these fuels as power sources.  
+
'''Sulfur dioxide''' is the [[chemical compound]] with the formula SO<sub>2</sub>.  It 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&Wiberg}}</ref> This is one of the causes for concern over the environmental impact of the use of these fuels as power sources.  
  
 
==Preparation==
 
==Preparation==
 
Sulfur dioxide can be prepared by burning [[sulfur]]:
 
Sulfur dioxide can be prepared by burning [[sulfur]]:
−
 
+
:S<sub>8</sub> + 8O<sub>2</sub> → 8SO<sub>2</sub>
−
:S<sub>8</sub> + 8 O<sub>2</sub> → 8 SO<sub>2</sub>
 
  
 
The combustion of [[hydrogen sulfide]] and organosulfur compounds proceeds similarly.
 
The combustion of [[hydrogen sulfide]] and organosulfur compounds proceeds similarly.
−
 
+
:2H<sub>2</sub>S + 3O<sub>2</sub> → 2H<sub>2</sub>O + 2 SO<sub>2</sub>
−
: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  [[pyrites]], [[sphalerite]] (zinc blende), and [[cinnabar]] (mercury sulfide) also releases SO<sub>2</sub>:
 
The roasting of sulfide ores such as  [[pyrites]], [[sphalerite]] (zinc blende), and [[cinnabar]] (mercury sulfide) also releases SO<sub>2</sub>:
−
 
+
:4FeS<sub>2</sub> + 11O<sub>2</sub> → 2Fe<sub>2</sub>O<sub>3</sub> + 8SO<sub>2</sub>
−
:4 FeS<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)
+
:2ZnS + 3O<sub>2</sub> → 2ZnO + 2SO<sub>2</sub>
−
 
+
:HgS + O<sub>2</sub> → Hg + SO<sub>2</sub>
−
:2 ZnS (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 [[calcium silicate]] [[cement]]: [[Calcium sulfate|CaSO<sub>4</sub>]] is heated with [[coke (fuel)|coke]] and sand in this process:
 
Sulfur dioxide is a by-product in the manufacture of [[calcium silicate]] [[cement]]: [[Calcium sulfate|CaSO<sub>4</sub>]] is heated with [[coke (fuel)|coke]] and sand in this process:
−
:2 CaSO<sub>4</sub> (s) + 2 SiO<sub>2</sub> (s) + C (s) → 2 CaSiO<sub>3</sub> (s) + 2 SO<sub>2</sub> (g) + CO<sub>2</sub> (g)
+
:2CaSO<sub>4</sub> + 2SiO<sub>2</sub> + C → 2CaSiO<sub>3</sub> + 2SO<sub>2</sub> + CO<sub>2</sub>
  
 
Action of hot [[sulfuric acid]] on copper [[swarf|turnings]] produces sulfur dioxide.
 
Action of hot [[sulfuric acid]] on copper [[swarf|turnings]] produces sulfur dioxide.
 +
:Cu + 2H<sub>2</sub>SO<sub>4</sub> → CuSO<sub>4</sub> + SO<sub>2</sub> + 2H<sub>2</sub>O
  
−
:Cu (s) + 2 H<sub>2</sub>SO<sub>4</sub> (aq) → CuSO<sub>4</sub> (aq) + SO<sub>2</sub> (g) + 2 H<sub>2</sub>O (l)
+
It can also be prepared from [[sodium metabisulfite]]:
−
 
+
:H<sub>2</sub>SO<sub>4</sub> +  Na<sub>2</sub>S<sub>2</sub>O<sub>5</sub> → 2SO<sub>2</sub> + Na<sub>2</sub>SO<sub>4</sub> + H<sub>2</sub>O
−
It can also be prepared with [[sodium metabisulfite]]:
 
−
 
 
−
:H<sub>2</sub>SO<sub>4</sub> (aq) +  Na<sub>2</sub>S<sub>2</sub>O<sub>5</sub> (aq) → 2 SO<sub>2</sub> (g) + Na<sub>2</sub>SO<sub>4</sub> (s) + H<sub>2</sub>O (l)
 
−
 
 
−
This is an exothermic reaction.
 
  
 
==Structure and bonding==
 
==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.
+
SO<sub>2</sub> is a bent molecule with ''C''<sub>2v</sub> [[symmetry point group|symmetry]].  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.
 
[[Image:Sulfur-dioxide-resonance-2D.png|500px|three [[resonance (chemistry)|resonance structures]] of sulfur dioxide]]
 
[[Image:Sulfur-dioxide-resonance-2D.png|500px|three [[resonance (chemistry)|resonance structures]] of sulfur dioxide]]
  
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==Reactions==
 
==Reactions==
 
Treatment of basic solutions with sulfur dioxide affords [[sulfite]] salts:
 
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
+
:SO<sub>2</sub> + 2NaOH → 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 to give the sulfuryl halides, such as [[sulfuryl chloride]]:
 
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]]:
−
 
 
:SO<sub>2</sub> + Cl<sub>2</sub> → SO<sub>2</sub>Cl<sub>2</sub>
 
:SO<sub>2</sub> + Cl<sub>2</sub> → 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:
 
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
+
:SO<sub>2</sub> + 2H<sub>2</sub>S → 3S + 2H<sub>2</sub>O
  
−
Sulfur dioxide can bind to metal ions as a [[ligand]], typically where the transition metal is in oxidation state 0 or +1.<ref name = "Greenwood">{{Greenwood&Earnshaw1st|pages=824–32}}</ref>  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 [[hapticity|η]]<sup>1</sup>.<ref name = "Greenwood"/><!--these are pretty specialized and probably should be in a separate article:
+
Sulfur dioxide can bind to metal ions as a [[ligand]], typically where the transition metal is in oxidation state 0 or +1.<ref name = "Greenwood">{{Greenwood&Earnshaw1st|pages=824–32}}.</ref>  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 [[hapticity|η]]<sup>1</sup>.<ref name = "Greenwood"/><!--these are pretty specialized and probably should be in a separate article:
 
*O donation  [[hapticity|η]]<sup>1</sup>
 
*O donation  [[hapticity|η]]<sup>1</sup>
 
*π donation- side  on [[hapticity|η]]<sup>2</sup>
 
*π donation- side  on [[hapticity|η]]<sup>2</sup>
Line 111: Line 98:
  
 
===As a preservative===
 
===As a preservative===
−
Sulfur dioxide is sometimes used as a [[preservative]] for dried [[apricot]]s and other [[dried fruit]]s owing to its [[antimicrobial]] properties, and it is sometimes called [[E number|E220]] when used in this way.  As a preservative, it maintains the appearance of the fruit and prevents [[Decomposition|rotting]]. Its presence gives the fruit a distinctive chemical [[taste]].{{Fact|date=June 2009}}
+
Sulfur dioxide is sometimes used as a [[preservative]] for dried [[apricot]]s and other [[dried fruit]]s owing to its [[antimicrobial]] properties, and it is sometimes called [[E number|E220]] when used in this way.  As a preservative, it maintains the appearance of the fruit and prevents [[Decomposition|rotting]].
  
 
===In winemaking===
 
===In winemaking===
−
Sulfur dioxide is an 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 an antibiotic and 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 10 ppm 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 350 ppm in the 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.{{Fact|May 2009|date=May 2009}}
+
Sulfur dioxide is an important compound in winemaking, and is designated as parts per million in wine, [[E number]]: E220.<ref>{{citation | url = http://www.food.gov.uk/safereating/chemsafe/additivesbranch/enumberlist#h_3 | webpage = Current EU approved additives and their E Numbers | publisher = UK Food Standards Agency}}.</ref>  It is present even in so-called unsulfurated wine at concentrations of up to 10&nbsp;milligrams per litre.<ref>{{citation | url = http://www.morethanorganic.com/sulphur-in-the-bottle | webpage = Sulphites in wine | website = MoreThanOrganic.com}}.</ref>  It serves as an antibiotic and antioxidant, protecting wine from spoilage by bacteria and oxidation.  
 +
 
 +
The presence of sulfur dioxide (at more than 10&nbsp;ppm) is indicated by the words "contains sulfites" found on wine labels in the US and EU. The upper limit of SO<sub>2</sub> allowed in wine is 350&nbsp;ppm in the US; in the EU it is 160&nbsp;ppm for red wines and 210&nbsp;ppm for white and 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 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 and sanitize equipment. Compounds of ozone (O<sub>3</sub>) are now used extensively as cleaning products in wineries due to their efficiency, and because these compounds do not affect the wine or equipment.
 
SO<sub>2</sub> is also a very important element in winery sanitation.  Wineries and equipment must be kept 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 and sanitize equipment. Compounds of ozone (O<sub>3</sub>) are now used extensively as cleaning products in wineries due to their efficiency, and because these compounds do not affect the wine or equipment.
  
 
===As a reducing agent===
 
===As a reducing agent===
−
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.  In municipal wastewater treatment sulfur dioxide is used to treat chlorinated wastewater prior to release.  Sulfur dioxide reduces free and combined chlorine to [[chloride]].<ref>Tchobanoglous, George. Wastewater Engineering. 3rd ed. New York: Mc Graw Hill, 1979.</ref>
+
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.  In municipal wastewater treatment sulfur dioxide is used to treat chlorinated wastewater prior to release.  Sulfur dioxide reduces free and combined chlorine to [[chloride]].<ref>{{citation | last = Tchobanoglous | first = George | title = Wastewater Engineering | edition = 3rd | location = New York | publisher = McGraw-Hill | year = 1979}}.</ref>
  
 
===Biochemical and biomedical roles===
 
===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]].
+
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===
 
===As a refrigerant===
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===As a reagent and solvent in the laboratory===
 
===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 salt]]s with sulfur dioxide and cuprous chloride affords the corresponding aryl sulfonyl chloride.<ref>{{OrgSynth | author = R. V. Hoffman | title = m-Trifluoromethylbenzenesulfonyl Chloride | collvol = 7 | collvolpages = 508 | year = 1990| prep = CV7P0508}}</ref>
+
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 and cuprous chloride affords the corresponding aryl sulfonyl chloride.<ref>{{OrgSynth | first = R. V. | last = Hoffman | title = m-Trifluoromethylbenzenesulfonyl Chloride | prep = CV7P0508 | volume = 60 | pages = 121 | year = 1981 | collvol = 7 | collvolpages = 508}}.</ref>
  
 
==Emissions==
 
==Emissions==
 
[[Image:Sulfur dioxide emissions from the Halema`uma`u vent, glows at night.jpg|thumb|Sulfur dioxide emissions from the [[Halemaumau Crater|Halema`uma`u vent]], glows at night]]
 
[[Image:Sulfur dioxide emissions from the Halema`uma`u vent, glows at night.jpg|thumb|Sulfur dioxide emissions from the [[Halemaumau Crater|Halema`uma`u vent]], glows at night]]
−
According to the [[United States Environmental Protection Agency]] (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]]:
+
According to the [[United States Environmental Protection Agency]] (EPA), the following amount of sulfur dioxide was released in the U.S. per year, measured in thousands of [[short tons]]:<ref>{{citation | url = http://www.epa.gov/air/airtrends/sulfur.html | webpage = National Trends in Sulfur Dioxide Levels | publisher = U.S. Environmental Protection Agency}}.</ref>
−
{| border=1
+
 
 +
{| class="wikitable"
 
|*'''1999'''
 
|*'''1999'''
 
|18,867
 
|18,867
Line 162: Line 152:
 
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.
 
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://www.physorg.com/news78159898.html China has its worst spell of acid rain], [[United Press International]].</ref>
+
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>{{citation | url = http://www.physorg.com/news78159898.html | webpage = China has its worst spell of acid rain | publisher = United Press International}}.</ref>
−
 
 
−
<!--anecdotal and notable?: [[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==
 
==Temperature dependence of aqueous solubility==
  
−
{|border=1
+
{| class="wikitable"
 
|22 g/100ml (0 °C)
 
|22 g/100ml (0 °C)
 
|15 g/100ml (10 °C)
 
|15 g/100ml (10 °C)
Line 190: Line 178:
  
 
==Safety==
 
==Safety==
−
Sulfur dioxide is an allergen to which some consumers are sensitive.  SO<sub>2</sub> is associated with increased respiratory symptoms and disease, difficulty in breathing, and       premature death.<ref>[http://www.epa.gov/air/urbanair/so2/hlth1.html][[United States Environmental Protection Agency‎|U.S. Environmental Protection Agency‎]]</ref>
+
Sulfur dioxide is an allergen to which some consumers are sensitive.  SO<sub>2</sub> is associated with increased respiratory symptoms and disease, difficulty in breathing, and premature death.<ref>{{citation | url = http://www.epa.gov/air/urbanair/so2/hlth1.html | publisher = U.S. Environmental Protection Agency‎}}.</ref>
  
 
==References==
 
==References==

Revision as of 18:47, 24 August 2009

Sulfur dioxide
Sulfur-dioxide-2D.svg
Sulfur-dioxide-3D-vdW.png
Other names Sulfur(IV) oxide
Sulfurous anhydride
Identifiers
CAS number [7446-09-5]
EC number 231-195-2
RTECS WS4550000
ChemSpider 1087
Properties[1]
Chemical formula SO2
Molar mass 64.064 g/mol
Appearance colorless gas
Density 2.551 g/L (gas)
1.46 g/cm3 (liquid, −10 °C)
Melting point

-75.5 °C, 198 K, -104 °F

Boiling point

-10.0 °C, 263 K, 14 °F

Solubility in water 9.4 g/100 mL (25 °C)
Solubility very soluble in acetone, methyl isobutyl ketone, acetic acid, alcohol
soluble in sulfuric acid
Acidity (pKa) 1.81
Viscosity 0.403 cP (0 °C)
Structure
Molecular geometry Bent, C2v
O–S–O = 119º
Dipole moment 1.62 D
Hazards
Material safety data sheet (MSDS) ICSC 0074
EU index number 016-011-00-9
Flash point non-flammable
LD50 3000 ppm (30 min inhaled, mouse)
Related compounds
Other cations Ozone
Selenium dioxide
Tellurium dioxide
Other sulfur oxides Sulfur monoxide
Sulfur trioxide
Other compounds sulfurous acid
Except where noted otherwise, data are given for materials in their standard state (at 25 °C, 100 kPa)

Sulfur dioxide is the chemical compound with the formula SO2. It is produced by volcanoes and in various industrial processes. Since coal and petroleum often contain sulfur compounds, their combustion generates sulfur dioxide. Further oxidation of SO2, usually in the presence of a catalyst such as NO2, forms H2SO4, and thus acid rain.[2] 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:

S8 + 8O2 → 8SO2

The combustion of hydrogen sulfide and organosulfur compounds proceeds similarly.

2H2S + 3O2 → 2H2O + 2 SO2

The roasting of sulfide ores such as pyrites, sphalerite (zinc blende), and cinnabar (mercury sulfide) also releases SO2:

4FeS2 + 11O2 → 2Fe2O3 + 8SO2
2ZnS + 3O2 → 2ZnO + 2SO2
HgS + O2 → Hg + SO2

Sulfur dioxide is a by-product in the manufacture of calcium silicate cement: CaSO4 is heated with coke and sand in this process:

2CaSO4 + 2SiO2 + C → 2CaSiO3 + 2SO2 + CO2

Action of hot sulfuric acid on copper turnings produces sulfur dioxide.

Cu + 2H2SO4 → CuSO4 + SO2 + 2H2O

It can also be prepared from sodium metabisulfite:

H2SO4 + Na2S2O5 → 2SO2 + Na2SO4 + H2O

Structure and bonding

SO2 is a bent molecule with C2v symmetry. In terms of electron-counting formalisms, the sulfur atom has an oxidation state of +4, a formal charge of 0, and 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. three resonance structures of sulfur dioxide

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]

Reactions

Treatment of basic solutions with sulfur dioxide affords sulfite salts:

SO2 + 2NaOH → 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

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:

SO2 + 2H2S → 3S + 2H2O

Sulfur dioxide can bind to metal ions as a ligand, typically where the transition metal is in oxidation state 0 or +1.[4] 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.[4]

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.[5] It is present even in so-called unsulfurated wine at concentrations of up to 10 milligrams per litre.[6] It serves as an antibiotic and antioxidant, protecting wine from spoilage by bacteria and oxidation.

The presence of sulfur dioxide (at more than 10 ppm) is indicated by the words "contains sulfites" found on wine labels in the US and EU. The upper limit of SO2 allowed in wine is 350 ppm in the US; 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 over 50ppm, SO2 becomes evident in the nose and taste of wine.

SO2 is also a very important element in winery sanitation. Wineries and equipment must be kept clean, and because bleach cannot be used in a winery, 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 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.[7]

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 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.[8]

Emissions

Sulfur dioxide emissions from the Halema`uma`u vent, glows at night

According to the United States Environmental Protection Agency (EPA), the following amount of sulfur dioxide was released in the U.S. per year, measured in thousands of short tons:[9]

*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 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 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, gypsum. Most gypsum sold in Europe comes from flue gas desulfurization.

New fuel additive catalysts, such as 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.[10]

Temperature dependence of aqueous solubility

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 SO2. Solubility of gas in a liquid depends on the gas partial pressure according to Henry's law. The solubility is given for "pure water", i.e., water that contains only SO2 in the amount at equilibrium with the gas phase. This "pure water" is going to be acidic. The solubility of SO2 in neutral (or alkaline) water is generally going to be higher because of the pH-dependent speciation of SO2 in the solution with the production of bisulfite and some sulfite ions.

Safety

Sulfur dioxide is an allergen to which some consumers are sensitive. SO2 is associated with increased respiratory symptoms and disease, difficulty in breathing, and premature death.[11]

References

  1. ↑ Patnaik, Pradyot Handbook of Inorganic Chemicals; McGraw-Hill, 2002. ISBN 0070494398.
  2. ↑ Holleman, A. F.; Wiberg, E. Inorganic Chemistry; Academic Press: San Diego, 2001. ISBN 0-12-352651-5
  3. ↑ Greenwood, N. N.; Earnshaw, A. Chemistry of the Elements, 2nd ed.; Butterworth-Heinemann: Oxford, 1997; p 700. ISBN 0-7506-3365-4
  4. ↑ 4.0 4.1 Greenwood, Norman N.; Earnshaw, A. Chemistry of the Elements; Pergamon: Oxford, 1984; pp 824–32. ISBN 0-08-022057-6.
  5. ↑ Current EU approved additives and their E Numbers, <http://www.food.gov.uk/safereating/chemsafe/additivesbranch/enumberlist#h_3>, UK Food Standards Agency.
  6. ↑ Sulphites in wine, <http://www.morethanorganic.com/sulphur-in-the-bottle>, MoreThanOrganic.com; .
  7. ↑ Tchobanoglous, George Wastewater Engineering, 3rd ed.; McGraw-Hill: New York, 1979.
  8. ↑ Hoffman, R. V. m-Trifluoromethylbenzenesulfonyl Chloride. Org. Synth. 1981, 60, 121, <http://www.orgsyn.org/orgsyn/orgsyn/prepContent.asp?prep=CV7P0508>; Coll. Vol., 7, 508.
  9. ↑ National Trends in Sulfur Dioxide Levels, <http://www.epa.gov/air/airtrends/sulfur.html>, U.S. Environmental Protection Agency.
  10. ↑ China has its worst spell of acid rain, <http://www.physorg.com/news78159898.html>, United Press International.
  11. ↑ , <http://www.epa.gov/air/urbanair/so2/hlth1.html>.

Further reading

External links

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