| Revision as of 04:35, 17 February 2012 editBeetstra (talk | contribs)Edit filter managers, Administrators172,031 edits Saving copy of the {{chembox}} taken from revid 477309810 of page Sulfur_dioxide for the Chem/Drugbox validation project (updated: 'ChEMBL'). |
Latest revision as of 22:35, 20 December 2024 edit EMsmile (talk | contribs)Event coordinators, Extended confirmed users59,919 edits →Air pollution: improved captionTag: Visual edit |
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{{Short description|Chemical compound of sulfur and oxygen}} |
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{{ambox | text = This page contains a copy of the infobox ({{tl|chembox}}) taken from revid of page ] with values updated to verified values.}} |
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{{cs1 config|name-list-style=vanc|display-authors=etal}} |
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{{Use American English|date=August 2020}} |
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{{Use mdy dates|date=August 2024}} |
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{{Chembox |
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{{Chembox |
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|Watchedfields = changed |
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| verifiedrevid = 477162682 |
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|verifiedrevid = 477313199 |
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| ImageFile = Sulfur-dioxide-2D.svg |
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|ImageFileL1 = Sulfur-dioxide-2D.svg |
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| ImageFile_Ref = {{chemboximage|correct|??}} |
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|ImageFileL1_Ref = {{chemboximage|correct|??}} |
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| ImageSize = 160 |
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|ImageSizeL1 = 160 |
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| ImageName = Skeletal formula sulfur dioxide with assorted dimensions |
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|ImageNameL1 = Skeletal formula sulfur dioxide with assorted dimensions |
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| ImageFile1 = Sulfur-dioxide-3D-vdW.png |
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|ImageFile2 = Sulfur-dioxide-3D-vdW.png |
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| ImageFile1_Ref = {{chemboximage|correct|??}} |
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|ImageFile2_Ref = {{chemboximage|correct|??}} |
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| ImageSize1 = 140 |
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|ImageSize2 = 120 |
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| ImageName1 = Spacefill model of sulfur dioxide |
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| IUPACName = Sulfur dioxide |
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|ImageName2 = Spacefill model of sulfur dioxide |
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|ImageFileR1 = Sulfur-dioxide-ve-B-2D.png |
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| OtherNames = Sulfurous anhydride<br /> |
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|ImageFileR1_Ref = {{chemboximage|correct|??}} |
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Sulfur(IV) oxide |
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|ImageSizeR1 = 160 |
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| Section1 = {{Chembox Identifiers |
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|ImageNameR1 = The Lewis structure of sulfur dioxide (SO2), showing unshared electron pairs. |
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| CASNo = 7446-09-5 |
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|IUPACName = Sulfur dioxide |
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| CASNo_Ref = {{cascite|correct|CAS}} |
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|OtherNames = {{Unbulleted list|Sulfurous anhydride|Sulfur(IV) oxide}} |
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| PubChem = 1119 |
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|Section1={{Chembox Identifiers |
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| PubChem_Ref = {{Pubchemcite|correct|pubchem}} |
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|CASNo = 7446-09-5 |
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| ChemSpiderID = 1087 |
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| ChemSpiderID_Ref = {{chemspidercite|correct|chemspider}} |
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|CASNo_Ref = {{cascite|correct|CAS}} |
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|PubChem = 1119 |
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| UNII = 0UZA3422Q4 |
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|ChemSpiderID = 1087 |
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| UNII_Ref = {{fdacite|correct|FDA}} |
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|ChemSpiderID_Ref = {{chemspidercite|correct|chemspider}} |
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| EINECS = 231-195-2 |
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|UNII = 0UZA3422Q4 |
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| UNNumber = 1079, 2037 |
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|UNII_Ref = {{fdacite|correct|FDA}} |
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| KEGG = D05961 |
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|EINECS = 231-195-2 |
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| KEGG_Ref = {{keggcite|correct|kegg}} |
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|UNNumber = 1079, 2037 |
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| MeSHName = Sulfur+dioxide |
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|KEGG = D05961 |
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| ChEBI_Ref = {{ebicite|correct|EBI}} |
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|KEGG_Ref = {{keggcite|correct|kegg}} |
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| ChEBI = 18422 |
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|MeSHName = Sulfur+dioxide |
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| ChEMBL = <!-- blanked - oldvalue: 1235997 --> |
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| ChEMBL_Ref = {{ebicite|changed|EBI}} |
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|ChEBI_Ref = {{ebicite|correct|EBI}} |
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|ChEBI = 18422 |
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| RTECS = WS4550000 |
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|ChEMBL = 1235997 |
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| Beilstein = 3535237 |
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|ChEMBL_Ref = {{ebicite|changed|EBI}} |
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| Gmelin = 1443 |
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|RTECS = WS4550000 |
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| SMILES = O=S=O |
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|Beilstein = 3535237 |
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| StdInChI = 1S/O2S/c1-3-2 |
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|Gmelin = 1443 |
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| StdInChI_Ref = {{stdinchicite|correct|chemspider}} |
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|SMILES = O=S=O |
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| InChI = 1/O2S/c1-3-2 |
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|StdInChI = 1S/O2S/c1-3-2 |
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| StdInChIKey = RAHZWNYVWXNFOC-UHFFFAOYSA-N |
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| StdInChIKey_Ref = {{stdinchicite|correct|chemspider}} |
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|StdInChI_Ref = {{stdinchicite|correct|chemspider}} |
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|InChI = 1/O2S/c1-3-2 |
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| InChIKey = RAHZWNYVWXNFOC-UHFFFAOYAT |
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|StdInChIKey = RAHZWNYVWXNFOC-UHFFFAOYSA-N |
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|StdInChIKey_Ref = {{stdinchicite|correct|chemspider}} |
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|InChIKey = RAHZWNYVWXNFOC-UHFFFAOYAT |
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}} |
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}} |
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| Section2 = {{Chembox Properties |
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|Section2={{Chembox Properties |
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| Formula = {{Chem|SO|2}} |
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|Formula = {{Chem|SO|2}} |
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| MolarMass = 64.066 g mol<sup>−1</sup> |
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|MolarMass = 64.066 g/mol |
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|Appearance = Colorless and pungent gas |
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| ExactMass = 63.961899934 g mol<sup>−1</sup> |
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|Density = 2.619 kg m<sup>−3</sup><ref>{{cite web |title=Sulfur Dioxide |url=https://pubchem.ncbi.nlm.nih.gov/compound/1119 |access-date=2024-03-22 |archive-date=2023-09-24 |archive-url=https://web.archive.org/web/20230924183723/https://pubchem.ncbi.nlm.nih.gov/compound/1119 |url-status=dead }}, U.S. National Library of Medicine</ref> |
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| Appearance = Colorless gas |
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|Solubility = 94 g/L<ref>{{RubberBible87th}}</ref><br />forms ] |
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| Density = 2.6288 kg m<sup>−3</sup> |
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|MeltingPtK = 201 |
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| Solubility = 94 g dm<sup>−3</sup><ref>{{RubberBible87th}}</ref> |
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|BoilingPtC = −10 |
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| MeltingPtK = 201 |
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|VaporPressure = 230 kPa at 10 °C; 330 kPa at 20 °C; 462 kPa at 30 °C; 630 kPa at 40 °C<ref>{{cite web |title=Hazardous Substances Data Bank |url=https://pubchem.ncbi.nlm.nih.gov/source/hsdb/228#section=Vapor-Pressure}}</ref> |
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| BoilingPtC = −10 |
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|pKa = ~1.81 |
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| VaporPressure = 237.2 kPa |
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| pKa = 1.81 |
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|pKb = ~12.19 |
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|Viscosity = 12.82 μPa·s<ref>{{cite journal |title=Correlation constants for chemical compounds|journal=Chemical Engineering |date=1976|last1=Miller|first1=J.W. Jr.|last2=Shah|first2=P.N.|last3=Yaws|first3=C.L.|volume=83|issue=25|pages=153–180|issn=0009-2460}}</ref> |
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| pKb = 12.19 |
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|Odor = Pungent; similar to a just-struck match<ref> {{Webarchive|url=https://web.archive.org/web/20191230172725/https://toxtown.nlm.nih.gov/text_version/chemicals.php?id=29 |date=December 30, 2019 }}, U.S. National Library of Medicine</ref> |
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| Viscosity = 0.403 cP (at 0 °C) |
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|MagSus = −18.2·10<sup>−6</sup> cm<sup>3</sup>/mol |
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}} |
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}} |
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| Section3 = {{Chembox Structure |
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|Section3={{Chembox Structure |
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| SpaceGroup = ''C''<sub>2''v''</sub> |
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|PointGroup = ''C''<sub>2''v''</sub> |
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| Coordination = Digonal |
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|Coordination = Digonal |
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| MolShape = Dihedral |
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|MolShape = Dihedral |
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| Dipole = 1.62 D |
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|Dipole = 1.62 D |
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}} |
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}} |
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| Section4 = {{Chembox Thermochemistry |
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|Section4={{Chembox Thermochemistry |
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| DeltaHf = -296.81 kJ mol<sup>−1</sup> |
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|DeltaHf = −296.81 kJ mol<sup>−1</sup> |
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| Entropy = 248.223 J K<sup>−1</sup> mol<sup>−1</sup> |
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|Entropy = 248.223 J K<sup>−1</sup> mol<sup>−1</sup> |
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}} |
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}} |
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| Section5 = {{Chembox Hazards |
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|Section5={{Chembox Hazards |
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|GHSPictograms = {{GHS corrosion}} {{GHS skull and crossbones}} |
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| EUIndex = 016-011-00-9 |
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|GHSSignalWord = Danger |
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| EUClass = {{Hazchem T}} |
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|HPhrases = {{H-phrases|314|331}}<ref>{{Cite web|url=https://echa.europa.eu/information-on-chemicals/cl-inventory-database/-/discli/notification-details/115657/1409763|title=C&L Inventory}}</ref> |
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| RPhrases = {{R23}}, {{R34}}, {{R50}} |
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|NFPA-H = 3 |
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| SPhrases = {{S1/2}}, {{S9}}, {{S26}}, {{S36/37/39}}, {{S45}} |
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| NFPA-H = 3 |
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|NFPA-F = 0 |
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| NFPA-F = 0 |
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|NFPA-R = 0 |
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|IDLH = 100 ppm<ref name=PGCH>{{PGCH|0575}}</ref> |
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| NFPA-R = 0 |
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|REL = TWA 2 ppm (5 mg/m<sup>3</sup>) ST 5 ppm (13 mg/m<sup>3</sup>)<ref name=PGCH/> |
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| LD50 = 3000 ppm (30 min inhaled, mouse) |
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|PEL = TWA 5 ppm (13 mg/m<sup>3</sup>)<ref name=PGCH/> |
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|LC50 = 3000 ppm (mouse, 30 min)<br />2520 ppm (rat, 1 hr)<ref name=IDLH>{{IDLH|7446095|Sulfur dioxide}}</ref> |
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|LCLo = 993 ppm (rat, 20 min)<br />611 ppm (rat, 5 hr)<br />764 ppm (mouse, 20 min)<br />1000 ppm (human, 10 min)<br />3000 ppm (human, 5 min)<ref name=IDLH/> |
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}} |
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| Section8 = {{Chembox Related |
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|Section6={{Chembox Related |
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| Function = ] ]s |
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|OtherFunction_label = ] ]s |
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| OtherFunctn = ]<br/>] |
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|OtherFunction = ]<br />]<br />] |
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| OtherCpds = ]<br /> |
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|OtherCompounds = ]<br /> |
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'''Sulfur dioxide''' (]-recommended spelling) or '''sulphur dioxide''' (traditional ]) is the ] with the formula {{chem|]|]|2}}. It is a colorless gas with a pungent smell that is responsible for the odor of burnt matches. It is released naturally by ] and is produced as a by-product of copper extraction and the burning of ]-] fossil fuels.<ref name=Greenwood/> |
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Sulfur dioxide is somewhat toxic to humans, although only when inhaled in relatively large quantities for a period of several minutes or more. It was known to medieval ] as "volatile spirit of sulfur".<ref name="Wothers-2019">{{Cite book |last=Wothers |first=Peter |url=https://books.google.com/books?id=8Cy7DwAAQBAJ&pg=PA69 |title=Antimony, Gold, and Jupiter's Wolf: How the Elements Were Named |date=2019 |publisher=Oxford University Press |isbn=978-0-19-965272-3 |language=en}}</ref> |
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== Structure and bonding == |
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SO<sub>2</sub> is a bent molecule with ''C''<sub>2v</sub> ]. |
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A ] approach considering just ''s'' and ''p'' orbitals would describe the bonding in terms of ] between two resonance structures. |
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] |
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The sulfur–oxygen bond has a ] of 1.5. There is support for this simple approach that does not invoke ''d'' orbital participation.<ref>{{cite journal |
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| title = Chemical bonding in oxofluorides of hypercoordinatesulfur |
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|author1=Cunningham, Terence P. |author2=Cooper, David L. |author3=Gerratt, Joseph |author4=Karadakov, Peter B. |author5=Raimondi, Mario | journal = Journal of the Chemical Society, Faraday Transactions |
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| year = 1997 |
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| volume = 93 |
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| issue = 13 |
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| pages = 2247–2254 |
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| doi = 10.1039/A700708F |
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}}</ref> |
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In terms of ] formalism, the sulfur atom has an ] of +4 and a ] of +1. |
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==Occurrence== |
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] |
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Sulfur dioxide is found on Earth and exists in very small concentrations in the atmosphere at about 15 ].<ref>{{Cite web |last=US EPA |first=OAR |date=May 4, 2016 |title=Sulfur Dioxide Trends |url=https://www.epa.gov/air-trends/sulfur-dioxide-trends |access-date=2023-02-16 |website=www.epa.gov |language=en}}</ref> |
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On other planets, sulfur dioxide can be found in various concentrations, the most significant being the ], where it is the third-most abundant atmospheric gas at 150 ppm. There, it reacts with water to form clouds of Sulfurous acid (SO2 + H2O ⇌ HSO−3+ H+), and is a key component of the planet's global atmospheric ] and contributes to ].<ref name="MarcqBertaux2012">{{cite journal|last1=Marcq|first1=Emmanuel|last2=Bertaux|first2=Jean-Loup|last3=Montmessin|first3=Franck|last4=Belyaev|first4=Denis|title=Variations of sulphur dioxide at the cloud top of Venus's dynamic atmosphere|journal=Nature Geoscience|volume=6|pages=25–28|year=2012|issue=1 |issn=1752-0894|doi=10.1038/ngeo1650|bibcode=2013NatGe...6...25M|s2cid=59323909}}</ref> It has been implicated as a key agent in the warming of early ], with estimates of concentrations in the lower atmosphere as high as 100 ppm,<ref name="HalevyZuber2007">{{cite journal|last1=Halevy|first1=I.|last2=Zuber|first2=M. T.|last3=Schrag|first3=D. P.|title=A Sulfur Dioxide Climate Feedback on Early Mars|journal=Science|volume=318|issue=5858|year=2007|pages=1903–1907|issn=0036-8075|doi=10.1126/science.1147039|pmid=18096802|bibcode=2007Sci...318.1903H|s2cid=7246517}}</ref> though it only exists in trace amounts. On both Venus and Mars, as on Earth, its primary source is thought to be volcanic. The ], a natural satellite of ], is 90% sulfur dioxide<ref name="IobookChap10">{{cite book |last=Lellouch |first=E. |editor=Lopes, R. M. C. |editor1-link=Rosaly Lopes |editor2=Spencer, J. R. |title=Io after Galileo |date=2007 |publisher=Springer-Praxis |isbn=978-3-540-34681-4 |pages=231–264 |chapter=Io's atmosphere }}</ref> and trace amounts are thought to also exist in the ]. The ] has observed the presence of sulfur dioxide on the ] ], where it is formed through ] in the planet's atmosphere.<ref>{{cite web | url=https://phys.org/news/2022-11-james-webb-space-telescope-reveals.html?adlt=strict&toWww=1&redig=BF6DD536430F43AF981737C5EEABD064 | title=James Webb Space Telescope reveals an exoplanet atmosphere as never seen before }}</ref> |
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As an ice, it is thought to exist in abundance on the ]—as subliming ice or frost on the trailing hemisphere of ],<ref name="CruikshankHowell1985">{{cite book |doi=10.1007/978-94-009-5418-2_55 |chapter=Sulfur Dioxide Ice on IO |title=ICES in the Solar System |year=1985 |last1=Cruikshank |first1=D. P. |last2=Howell |first2=R. R. |last3=Geballe |first3=T. R. |last4=Fanale |first4=F. P. |pages=805–815 |isbn=978-94-010-8891-6 }}</ref> and in the crust and mantle of ], ], and ], possibly also in liquid form and readily reacting with water.<ref>. Jpl.nasa.gov (October 4, 2010). Retrieved on September 24, 2013.</ref> |
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==Production== |
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Sulfur dioxide is primarily produced for ] manufacture (see ], but other processes predated that at least since 16th century<ref name="Wothers-2019" />). In the United States in 1979, 23.6 million metric tons (26 million U.S. short tons) of sulfur dioxide were used in this way, compared with 150,000 metric tons (165,347 U.S. short tons) used for other purposes. Most sulfur dioxide is produced by the combustion of elemental ]. Some sulfur dioxide is also produced by roasting ] and other ] ores in air.<ref name = Ullmann>{{Ullmann| author = Müller, Hermann | title = Sulfur Dioxide | doi = 10.1002/14356007.a25_569}}</ref> |
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]. A flow-chamber joined to a gas washing bottle (filled with a solution of ]) is being used. The product is sulfur dioxide (SO<sub>2</sub>) with some traces of ] (SO<sub>3</sub>). The "smoke" that exits the gas washing bottle is, in fact, a sulfuric acid fog generated in the reaction.]] |
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===Combustion routes=== |
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Sulfur dioxide is the product of the burning of ] or of burning materials that contain sulfur: |
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:{{chem2|S8}} + 8 {{chem2|O2}} → 8 {{chem2|SO2}}, ΔH = −297 kJ/mol |
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To aid combustion, liquified sulfur ({{convert|140|–|150|C|F}} is sprayed through an atomizing nozzle to generate fine drops of sulfur with a large surface area. The reaction is ], and the combustion produces temperatures of {{convert|1000|–|1600|C|F}}. The significant amount of heat produced is recovered by steam generation that can subsequently be converted to electricity.<ref name = Ullmann/> |
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The combustion of ] and organosulfur compounds proceeds similarly. For example: |
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:2 {{chem2|H2S}} + 3 {{chem2|O2}} → 2 {{chem2|SO2}} + 2 {{chem2|H2O}} |
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The ] of sulfide ores such as ], ], and ] (mercury sulfide) also releases SO<sub>2</sub>:<ref>Shriver, Atkins. Inorganic Chemistry, Fifth Edition. W. H. Freeman and Company; New York, 2010; p. 414.</ref> |
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:4 {{chem2|FeS2}} + 11 {{chem2|O2}} → 2 {{chem2|Fe2O3}} + 8 {{chem2|SO2}} |
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:2 {{chem2|ZnS}} + 3 {{chem2|O2}} → 2 {{chem2|ZnO}} + 2 {{chem2|SO2}} |
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:{{chem2|HgS + O2 -> Hg + SO2}} |
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:4 FeS + 7 {{chem2|O2}} → 2 {{chem2|Fe2O3}} + 4 {{chem2|SO2}} |
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A combination of these reactions is responsible for the largest source of sulfur dioxide, volcanic eruptions. These events can release millions of tons of SO<sub>2</sub>. |
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===Reduction of higher oxides=== |
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Sulfur dioxide can also be a byproduct in the manufacture of ] cement; ] is heated with ] and sand in this process: |
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:2 {{chem2|CaSO4}} + 2 {{chem2|SiO2}} + C → 2 {{chem2|CaSiO3}} + 2 {{chem2|SO2}} + {{chem2|CO2}} |
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Until the 1970s commercial quantities of sulfuric acid and cement were produced by this process in ], England. Upon being mixed with ] or ], and roasted, the sulfate liberated sulfur dioxide gas, used in sulfuric acid production, the reaction also produced calcium silicate, a precursor in cement production.<ref>. lakestay.co.uk (2007)</ref> |
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On a laboratory scale, the action of hot concentrated sulfuric acid on copper ] produces sulfur dioxide. |
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:Cu + 2 {{chem2|H2SO4}} → {{chem2|CuSO4 + SO2 + 2 H2O}} |
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Tin also reacts with concentrated sulfuric acid but it produces tin(II) sulfate which can later be pyrolyzed at 360 °C into tin dioxide and dry sulfur dioxide. |
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:Sn + {{chem2|H2SO4}} → {{chem2|SnSO4 + H2}} |
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:{{chem2|SnSO4}} → {{chem2|SnO2}} + {{chem2|SO2}} |
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===From sulfites=== |
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The reverse reaction occurs upon acidification: |
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:{{chem2|H+ + HSO3- -> SO2 + H2O}} |
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==Reactions== |
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Sulfites results by the action of aqueous base on sulfur dioxide: |
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:{{chem2|SO2 + 2 NaOH → Na2SO3 + H2O}} |
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Sulfur dioxide is a mild but useful ]. It is oxidized by halogens to give the sulfuryl halides, such as ]: |
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:{{chem2|SO2 + Cl2 → SO2Cl2}} |
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Sulfur dioxide is the ] in the ], which is conducted on a large scale in ]. Here, sulfur dioxide is reduced by hydrogen sulfide to give elemental sulfur: |
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:{{chem2|SO2 + 2 H2S → 3 S + 2 H2O}} |
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The sequential oxidation of sulfur dioxide followed by its hydration is used in the production of sulfuric acid. |
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:{{chem2|SO2}} + {{chem2|H2O}} + {{frac|1|2}} {{chem2|O2}} → {{chem2|H2SO4}} |
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Sulfur dioxide dissolves in water to give "]", which cannot be isolated and is instead an acidic solution of ], and possibly ], ions. |
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:{{chem2|SO2 + H2O ⇌ HSO3− + H+}}{{spaces|10}}''K''<sub>a</sub> = 1.54{{e|−2}}; p''K''<sub>a</sub> = 1.81 |
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===Laboratory reactions=== |
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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 ] solution, turning the solution from orange to green (Cr<sup>3+</sup> (aq)). It can also reduce ferric ions to ferrous.<ref name=Lucas >{{Cite web|url=http://publications.gc.ca/collections/collection_2017/rncan-nrcan/M34-20/M34-20-107-eng.pdf|title = Information archivée dans le Web}}</ref> |
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Sulfur dioxide can react with certain 1,3-]s in a ] to form cyclic ]s. This reaction is exploited on an industrial scale for the synthesis of ], which is an important solvent in the ]. |
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:] |
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Sulfur dioxide can bind to metal ions as a ] to form ]es, 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 ]<sup>1</sup>.<ref name=Greenwood>{{Greenwood&Earnshaw2nd}}</ref> As a η<sup>1</sup>-SO<sub>2</sub> (S-bonded planar) ligand sulfur dioxide functions as a Lewis base using the lone pair on S. SO<sub>2</sub> functions as a ] in its η<sup>1</sup>-SO<sub>2</sub> (S-bonded pyramidal) bonding mode with metals and in its 1:1 ] with Lewis bases such as ] and ]. When bonding to Lewis bases the ] of SO<sub>2</sub> are E<sub>A</sub> = 0.51 and E<sub>A</sub> = 1.56. |
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==Uses== |
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The overarching, dominant use of sulfur dioxide is in the production of ].<ref name = Ullmann/> |
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===Precursor to sulfuric acid=== |
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Sulfur dioxide is an intermediate in the production of sulfuric acid, being converted to ], and then to ], 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 ]. Several million tons are produced annually for this purpose. |
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===Food preservative=== |
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{{See also|Food preservation}} |
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Sulfur dioxide is sometimes used as a preservative for dried apricots, dried figs, and other dried fruits, owing to its ] properties and ability to prevent ],<ref>{{cite conference |last1=Zamboni |first1=Cibele B. |last2=Medeiros |first2=Ilca M. M. A. |last3=de Medeiros |first3=José A. G. |title=Analysis of Sulfur in Dried Fruits Using NAA |url=https://www.ipen.br/biblioteca/2011/inac/17204.pdf |conference=2011 International Nuclear Atlantic Conference – INAC 2011 |isbn=978-85-99141-03-8 |date=October 2011 |access-date=2020-06-04 |archive-date=2020-06-04 |archive-url=https://web.archive.org/web/20200604193519/https://www.ipen.br/biblioteca/2011/inac/17204.pdf |url-status=dead }}</ref> and is called ]220<ref>, The Food Standards Agency website.</ref> when used in this way in Europe. As a preservative, it maintains the colorful appearance of the fruit and prevents ]. Historically, ] was "sulfured" as a preservative and also to lighten its color. Treatment of dried fruit was usually done outdoors, by igniting sublimed sulfur and burning in an enclosed space with the fruits.<ref name="University of Georgia">{{Citation |title=Preserving foods: Drying fruits and Vegetable |url=https://nchfp.uga.edu/publications/uga/uga_dry_fruit.pdf |publisher=University of Georgia cooperative extension service |access-date=2022-06-06 |archive-date=2022-09-27 |archive-url=https://web.archive.org/web/20220927163031/https://nchfp.uga.edu/publications/uga/uga_dry_fruit.pdf |url-status=dead }}</ref> Fruits may be sulfured by dipping them into an either ], ] or ].<ref name="University of Georgia" /> |
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==== Winemaking ==== |
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Sulfur dioxide was first used in ] by the Romans, when they discovered that burning sulfur candles inside empty wine vessels keeps them fresh and free from vinegar smell.<ref>{{cite web|url=http://www.practicalwinery.com/janfeb09/page1.htm|publisher=www.practicalwinery.com|date=February 1, 2009|title=Practical Winery & vineyard Journal Jan/Feb 2009|url-status=dead|archive-url=https://web.archive.org/web/20130928111625/http://www.practicalwinery.com/janfeb09/page1.htm|archive-date=2013-09-28}}</ref> |
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It is still an important compound in winemaking, and is measured in ] (''ppm'') in wine. It is present even in so-called unsulfurated wine at concentrations of up to 10 mg/L.<ref>, MoreThanOrganic.com.</ref> It serves as an ] and ], protecting wine from spoilage by bacteria and oxidation – a phenomenon that leads to the browning of the wine and a loss of cultivar specific flavors.<ref name="Jackson">Jackson, R.S. (2008) Wine science: principles and applications, Amsterdam; Boston: Elsevier/Academic Press</ref><ref name="Guerrero">{{cite journal | doi = 10.1016/j.tifs.2014.11.004| title = Demonstrating the efficiency of sulphur dioxide replacements in wine: A parameter review| journal = Trends in Food Science & Technology| volume = 42| pages = 27–43| year = 2015| last1 = Guerrero| first1 = Raúl F| last2 = Cantos-Villar| first2 = Emma| issue = 1}}</ref> Its antimicrobial action also helps minimize volatile acidity. Wines containing sulfur dioxide are typically labeled with "containing ]s". |
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Sulfur dioxide exists in wine in free and bound forms, and the combinations are referred to as total SO<sub>2</sub>. Binding, for instance to the carbonyl group of ], varies with the wine in question. The free form exists in equilibrium between molecular SO<sub>2</sub> (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) SO<sub>2</sub>, which is the active form, while at higher pH more SO<sub>2</sub> is found in the inactive sulfite and bisulfite forms. The molecular SO<sub>2</sub> is active as an antimicrobial and antioxidant, and this is also the form which may be perceived as a pungent odor at high levels. Wines with total 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 total SO<sub>2</sub> 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, SO<sub>2</sub> is mostly undetectable in wine, but at free SO<sub>2</sub> concentrations over 50 ppm, SO<sub>2</sub> becomes evident in the smell and taste of wine.{{Citation needed|date=May 2009}} |
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SO<sub>2</sub> 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 to the risk of ],<ref>. Purdue University</ref> a mixture of SO<sub>2</sub>, water, and citric acid is commonly used to clean and sanitize equipment. ] (O<sub>3</sub>) is now used extensively for sanitizing in wineries due to its efficacy, and because it does not affect the wine or most equipment.<ref> {{Webarchive|url=https://web.archive.org/web/20170912102459/https://www.practicalwinery.com/janfeb00/ozone.htm |date=September 12, 2017 }}, Practical Winery & Vineyard Journal.</ref> |
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===As a reducing agent=== |
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Sulfur dioxide is also a good ]. In the presence of water, sulfur dioxide is able to decolorize substances. Specifically, it is a useful reducing ] 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 ].<ref>{{cite book |last=Tchobanoglous |first=George |title=Wastewater Engineering |edition=3rd |location=New York |publisher=McGraw Hill |year=1979 |isbn=0-07-041677-X }}</ref> |
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Sulfur dioxide is fairly soluble in water, and by both IR and Raman spectroscopy; the hypothetical ], H<sub>2</sub>SO<sub>3</sub>, is not present to any extent. However, such solutions do show spectra of the hydrogen sulfite ion, HSO<sub>3</sub><sup>−</sup>, by reaction with water, and it is in fact the actual reducing agent present: |
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:SO<sub>2</sub> + H<sub>2</sub>O ⇌ HSO<sub>3</sub><sup>−</sup> + H<sup>+</sup> |
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===As a fumigant=== |
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In the beginning of the 20th century sulfur dioxide was used in ] as a fumigant to kill rats that carried the ] bacterium, which causes bubonic plague. The application was successful, and the application of this method was extended to other areas in South America. In Buenos Aires, where these apparatuses were known as ], but later also in Rio de Janeiro, New Orleans and San Francisco, the sulfur dioxide treatment machines were brought into the streets to enable extensive disinfection campaigns, with effective results.<ref>{{cite journal |last1=Engelmann |first1=Lukas |title=Fumigating the Hygienic Model City: Bubonic Plague and the Sulfurozador in Early-Twentieth-Century Buenos Aires |journal=Medical History |date=July 2018 |volume=62 |issue=3 |pages=360–382 |doi=10.1017/mdh.2018.37 |pmid=29886876 |pmc=6113751 }}</ref> |
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===Biochemical and biomedical roles=== |
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Sulfur dioxide 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.<ref>{{cite journal |last1=Liu |first1=D. |last2=Jin |first2=H. |last3=Tang |first3=C. |last4=Du |first4=J. |title=Sulfur Dioxide: a Novel Gaseous Signal in the Regulation of Cardiovascular Functions |journal=Mini-Reviews in Medicinal Chemistry |year=2010 |volume=10 |issue=11 |pages=1039–1045 |doi=10.2174/1389557511009011039 |pmid=20540708 }}</ref> Sulfur dioxide blocks nerve signals from the ] and abolishes the ]. |
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It is considered that endogenous sulfur dioxide plays a significant physiological role in regulating ] and ] function, and aberrant or deficient sulfur dioxide metabolism can contribute to several different cardiovascular diseases, such as ], ], ], and ].<ref>{{cite journal |last1=Tian |first1=Hong |title=Advances in the study on endogenous sulfur dioxide in the cardiovascular system |journal=Chinese Medical Journal |date=November 5, 2014 |volume=127 |issue=21 |pages=3803–3807 |doi=10.3760/cma.j.issn.0366-6999.20133031 |pmid=25382339 |s2cid=11924999 |url=https://journals.lww.com/25382339.pmid |doi-access=free }}</ref> |
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It was shown that in children with pulmonary arterial hypertension due to congenital heart diseases the level of ] is higher and the level of endogenous sulfur dioxide is lower than in normal control children. Moreover, these biochemical parameters strongly correlated to the severity of pulmonary arterial hypertension. Authors considered homocysteine to be one of useful biochemical markers of disease severity and sulfur dioxide metabolism to be one of potential therapeutic targets in those patients.<ref>{{cite journal|vauthors=Yang R, Yang Y, Dong X, Wu X, Wei Y |title=Correlation between endogenous sulfur dioxide and homocysteine in children with pulmonary arterial hypertension associated with congenital heart disease|language=zh|journal=Zhonghua Er Ke Za Zhi|date=Aug 2014|volume=52|issue=8|pages=625–629|pmid=25224243}}</ref> |
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Endogenous sulfur dioxide also has been shown to lower the ] rate of endothelial ] cells in blood vessels, via lowering the ] activity and activating ] and ].<ref>{{cite journal|vauthors=Liu D, Huang Y, Bu D, Liu AD, Holmberg L, Jia Y, Tang C, Du J, Jin H |title=Sulfur dioxide inhibits vascular smooth muscle cell proliferation via suppressing the Erk/MAP kinase pathway mediated by cAMP/PKA signaling|journal=Cell Death Dis.|date=May 2014|volume=5|issue=5|pages=e1251|doi=10.1038/cddis.2014.229|pmid=24853429|pmc=4047873}}</ref> Smooth muscle cell proliferation is one of important mechanisms of hypertensive remodeling of blood vessels and their ], so it is an important pathogenetic mechanism in arterial hypertension and atherosclerosis. |
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Endogenous sulfur dioxide in low concentrations causes endothelium-dependent ]. In higher concentrations it causes endothelium-independent vasodilation and has a negative inotropic effect on cardiac output function, thus effectively lowering blood pressure and myocardial oxygen consumption. The vasodilating and bronchodilating effects of sulfur dioxide are mediated via ATP-dependent ]s and L-type ("dihydropyridine") calcium channels. Endogenous sulfur dioxide is also a potent antiinflammatory, antioxidant and cytoprotective agent. It lowers blood pressure and slows hypertensive remodeling of blood vessels, especially thickening of their intima. It also regulates lipid metabolism.<ref>{{cite journal|vauthors=Wang XB, Jin HF, Tang CS, Du JB |title=The biological effect of endogenous sulfur dioxide in the cardiovascular system.|journal=Eur J Pharmacol|date=November 16, 2011|volume=670|issue=1|doi=10.1016/j.ejphar.2011.08.031|pmid=21925165|pages=1–6}}</ref> |
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Endogenous sulfur dioxide also diminishes myocardial damage, caused by ] ] hyperstimulation, and strengthens the myocardial antioxidant defense reserve.<ref>{{cite journal|vauthors=Liang Y, Liu D, Ochs T, Tang C, Chen S, Zhang S, Geng B, Jin H, Du J |title=Endogenous sulfur dioxide protects against isoproterenol-induced myocardial injury and increases myocardial antioxidant capacity in rats.|journal=Lab. Invest.|date=Jan 2011|volume=91|issue=1|pages=12–23|doi=10.1038/labinvest.2010.156|pmid=20733562|doi-access=free}}</ref> |
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===As a reagent and solvent in the laboratory=== |
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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 ]. Treatment of aryl ]s with sulfur dioxide and ] yields the corresponding aryl sulfonyl chloride, for example:<ref>{{OrgSynth | author = Hoffman, R. V. | title = m-Trifluoromethylbenzenesulfonyl Chloride | collvol = 7 | collvolpages = 508 | year = 1990| prep = CV7P0508}}</ref> |
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:] |
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As a result of its very low ], it is often used as a low-temperature solvent/diluent for superacids like ] (FSO<sub>3</sub>H/SbF<sub>5</sub>), allowing for highly reactive species like ''tert''-butyl cation to be observed spectroscopically at low temperature (though tertiary carbocations do react with SO<sub>2</sub> above about −30 °C, and even less reactive solvents like ] must be used at these higher temperatures).<ref>{{Cite journal|last1=Olah|first1=George A.|last2=Lukas|first2=Joachim.|date=August 1, 1967|title=Stable carbonium ions. XLVII. Alkylcarbonium ion formation from alkanes via hydride (alkide) ion abstraction in fluorosulfonic acid-antimony pentafluoride-sulfuryl chlorofluoride solution|journal=Journal of the American Chemical Society|volume=89|issue=18|pages=4739–4744|doi=10.1021/ja00994a030|issn=0002-7863}}</ref> |
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===As a refrigerant=== |
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Being easily condensed and possessing a high ], sulfur dioxide is a candidate material for refrigerants. Before the development of ]s, sulfur dioxide was used as a ] in ]. |
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===As an indicator of volcanic activity=== |
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Sulfur dioxide content in naturally-released geothermal gasses is measured by the ] as an indicator of possible volcanic activity.<ref>{{Cite web |date=n.d. |title=Volcanic gases |url=https://en.vedur.is/volcanoes/volcanic-hazards/volcanic-gases/ |website=Iceland Met Office}}</ref> |
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==Safety== |
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] volunteer tests for sulfur dioxide after the ].]] |
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===Ingestion=== |
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In the United States, the ] lists the two food preservatives, sulfur dioxide and ], as being safe for human consumption except for certain asthmatic individuals who may be sensitive to them, especially in large amounts.<ref>{{cite web | title = Center for Science in the Public Interest – Chemical Cuisine | url = http://www.cspinet.org/reports/chemcuisine.htm | access-date = March 17, 2010}}</ref> Symptoms of sensitivity to ] agents, including sulfur dioxide, manifest as potentially life-threatening trouble breathing within minutes of ingestion.<ref>{{cite web | title = California Department of Public Health: Food and Drug Branch: Sulfites | url = http://www.cdph.ca.gov/pubsforms/Guidelines/Documents/fdb%20Sulfites.pdf | access-date = September 27, 2013 | url-status = dead | archive-url = https://web.archive.org/web/20120723065412/http://www.cdph.ca.gov/pubsforms/Guidelines/Documents/fdb%20Sulfites.pdf | archive-date = July 23, 2012 }}</ref> Sulphites may also cause symptoms in non-asthmatic individuals, namely ], ], ], ], ] and diarrhea, and even life-threatening ].<ref name="pmid24834193">{{cite journal |vauthors=Vally H, Misso NL |title=Adverse reactions to the sulphite additives |journal=Gastroenterol Hepatol Bed Bench |volume=5 |issue=1 |pages=16–23 |date=2012 |pmid=24834193 |pmc=4017440 |doi= |url=}}</ref> |
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===Inhalation=== |
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Incidental exposure to sulfur dioxide is routine, e.g. the smoke from matches, coal, and sulfur-containing fuels like ]. Relative to other chemicals, it is only mildly toxic and requires high concentrations to be actively hazardous.<ref> U.S. Environmental Protection Agency</ref> However, its ubiquity makes it a major air pollutant with significant impacts on human health.<ref name="EPA">. ]</ref> |
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In 2008, the ] reduced the ] to 0.25 parts per million (ppm). In the US, the ] set the ] at 5 ppm (13 mg/m<sup>3</sup>) time-weighted average. Also in the US, ] set the ] at 100 ppm.<ref name=NIOSH>{{cite web |url=https://www.cdc.gov/niosh/npg/npgd0575.html |title=NIOSH Pocket Guide to Chemical Hazards }}</ref> In 2010, the ] "revised the primary SO<sub>2</sub> ] by establishing a new one-hour standard at a level of 75 ]. EPA revoked the two existing primary standards because they would not provide additional public health protection given a one-hour standard at 75 ppb."<ref name="EPA" /> |
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==Environmental role== |
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===Air pollution=== |
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]s on sulfate aerosol concentrations and chemical reactions in the atmosphere]] |
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Major ]s have an overwhelming effect on sulfate aerosol concentrations in the years when they occur: eruptions ranking 4 or greater on the ] inject {{chem2|SO2}} and water vapor directly into the ], where they react to create sulfate aerosol plumes.<ref name="nasa-aerosols">{{cite web |url=http://volcanoes.usgs.gov/hazards/gas/s02aerosols.php |title=Volcanic Sulfur Aerosols Affect Climate and the Earth's Ozone Layer |access-date=February 17, 2009 |publisher=United States Geological Survey |archive-date=November 14, 2015 |archive-url=https://web.archive.org/web/20151114184944/https://volcanoes.usgs.gov/hazards/gas/s02aerosols.php |url-status=dead }}</ref> Volcanic emissions vary significantly in composition, and have complex chemistry due to the presence of ash particulates and a wide variety of other elements in the plume. Only ] containing primarily ] magmas are responsible for these fluxes, as ] magma erupted in ] doesn't result in plumes which reach the stratosphere.<ref>{{cite journal |doi=10.1016/j.atmosenv.2004.06.017 |journal=Atmospheric Environment |volume=38 |issue=33 |year=2004 |pages=5637–5649 |title=Aerosol chemistry of emissions from three contrasting volcanoes in Italy |vauthors=Mathera TA, Oppenheimer AG, McGonigle A|bibcode=2004AtmEn..38.5637M }}</ref> However, before the ], dimethyl sulfide pathway was the largest contributor to sulfate aerosol concentrations in a more average year with no major volcanic activity. According to the ], published in 1990, volcanic emissions usually amounted to around 10 million tons in 1980s, while dimethyl sulfide amounted to 40 million tons. Yet, by that point, the global human-caused emissions of sulfur into the atmosphere became "at least as large" as ''all'' natural emissions of sulfur-containing compounds ''combined'': they were at less than 3 million tons per year in 1860, and then they increased to 15 million tons in 1900, 40 million tons in 1940 and about 80 millions in 1980. The same report noted that "in the industrialized regions of Europe and North America, anthropogenic emissions dominate over natural emissions by about a factor of ten or even more".<ref name="IPCC_FAR">IPCC, 1990: . In: . Cambridge University Press, Cambridge, United Kingdom and New York, NY, USA, pp. 31–34,</ref> In the eastern United States, sulfate particles were estimated to account for 25% or more of all air pollution.<ref name="EPAHealth" /> Exposure to sulfur dioxide emissions by coal power plants (coal PM<sub>2.5</sub>) in the US was associated with 2.1 times greater mortality risk than exposure to PM<sub>2.5</sub> from all sources.<ref name="science2023mortality">{{cite journal |last1=Henneman |first1=Lucas |last2=Choirat |first2=Christine |last3=Dedoussi |first3=Irene |last4=Dominici |first4=Francesca |last5=Roberts |first5=Jessica|last6=Zigler |first6=Corwin |title=Mortality risk from United States coal electricity generation |journal=] |date=November 24, 2023 |volume=382 |issue=6673 |pages=941–946|doi=10.1126/science.adf4915 |pmid=37995235 |pmc=10870829 |bibcode=2023Sci...382..941H |language=en}}</ref> |
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Meanwhile, the ] had much lower concentrations due to being much less densely populated, with an estimated 90% of the human population in the north. In the early 1990s, anthropogenic sulfur dominated in the ], where only 16% of annual sulfur emissions were natural, yet amounted for less than half of the emissions in the Southern Hemisphere.<ref>{{Cite journal|last1=Bates|first1=T. S.|last2=Lamb|first2=B. K.|last3=Guenther|first3=A.|last4=Dignon|first4=J.|last5=Stoiber|first5=R. E.|date=April 1992|title=Sulfur emissions to the atmosphere from natural sources|url=http://link.springer.com/10.1007/BF00115242|journal=Journal of Atmospheric Chemistry|language=en|volume=14|issue=1–4|pages=315–337|doi=10.1007/BF00115242 |bibcode=1992JAtC...14..315B |s2cid=55497518|issn=0167-7764}}</ref> |
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]]] |
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Such an increase in sulfate aerosol emissions had a variety of effects. At the time, the most visible one was ], caused by ] from clouds carrying high concentrations of sulfate aerosols in the ].<ref>{{Cite journal|last1=Burns|first1= Douglas A.|last2=Aherne|first2=Julian|last3=Gay|first3=David A.|last4=Lehmann|first4=Christopher M.~B.|title =Acid rain and its environmental effects: Recent scientific advances|journal = Atmospheric Environment|language=en|year = 2016|volume=146|pages = 1–4|doi = 10.1016/j.atmosenv.2016.10.019|bibcode= 2016AtmEn.146....1B|doi-access=free}}</ref> |
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At its peak, acid rain has eliminated ] and some other fish species and insect life from lakes and streams in geographically sensitive areas, such as ] in the United States.<ref name="EPASurface">{{Cite web|title=Effects of Acid Rain – Surface Waters and Aquatic Animals|url=http://www.epa.gov/acidrain/effects/surface_water.html|url-status=dead|archive-url=https://web.archive.org/web/20090514121649/http://www.epa.gov/acidrain/effects/surface_water.html|archive-date=May 14, 2009|website=US EPA}}</ref> Acid rain worsens ] function as some of its ] is lost and heavy metals like aluminium are mobilized (spread more easily) while essential nutrients and minerals such as ] can leach away because of the same. Ultimately, plants unable to tolerate lowered ] are killed, with montane forests being some of the worst-affected ]s due to their regular exposure to sulfate-carrying fog at high altitudes.<ref>{{Cite journal|last1=Rodhe|first1=Henning|last2=Dentener|first2=Frank|last3=Schulz|first3=Michael|date=October 1, 2002|title=The Global Distribution of Acidifying Wet Deposition|url=https://doi.org/10.1021/es020057g|journal=Environmental Science & Technology|volume=36|issue=20|pages=4382–4388|doi=10.1021/es020057g|pmid=12387412|bibcode=2002EnST...36.4382R|issn=0013-936X}}</ref><ref name="EPA: Forests">US EPA: {{webarchive |url=https://web.archive.org/web/20080726034352/http://www.epa.gov/acidrain/effects/forests.html |date=July 26, 2008 }}</ref><ref>{{cite journal|doi=10.1126/science.272.5259.244|url=http://www.esf.edu/efb/mitchell/Class%20Readings/Sci.272.244.246.pdf|title=Long-Term Effects of Acid Rain: Response and Recovery of a Forest Ecosystem|year=1996|last1=Likens|first1=G. E.|last2=Driscoll|first2=C. T.|last3=Buso|first3=D. C.|journal=Science|volume=272|issue=5259|page=244|bibcode=1996Sci...272..244L|s2cid=178546205|access-date=February 9, 2013|archive-date=December 24, 2012|archive-url=https://web.archive.org/web/20121224203613/http://www.esf.edu/efb/mitchell/Class%20Readings/Sci.272.244.246.pdf|url-status=live}}</ref><ref>{{Cite journal|last1=Larssen|first1=T.|last2=Carmichael|first2=G. R.|date=October 1, 2000|title=Acid rain and acidification in China: the importance of base cation deposition|url=http://www.sciencedirect.com/science/article/pii/S0269749199002791|journal=Environmental Pollution|language=en|volume=110|issue=1|pages=89–102|doi=10.1016/S0269-7491(99)00279-1|pmid=15092859|issn=0269-7491|access-date=April 22, 2020|archive-date=March 30, 2015|archive-url=https://web.archive.org/web/20150330041614/http://www.sciencedirect.com/science/article/pii/S0269749199002791|url-status=live}}</ref><ref>{{Cite journal|last1=Johnson|first1=Dale W.|last2=Turner|first2=John|last3=Kelly|first3=J. M.|date=1982|title=The effects of acid rain on forest nutrient status|journal=Water Resources Research|language=en|volume=18|issue=3|pages=449–461|doi=10.1029/WR018i003p00449|bibcode=1982WRR....18..449J|issn=1944-7973}}</ref> While acid rain was too dilute to affect human health directly, breathing smog or even any air with elevated sulfate concentrations is known to contribute to ] and ] conditions, including ] and ].<ref name="EPAHealth"> {{Webarchive|url=https://web.archive.org/web/20080118120242/http://www.epa.gov/acidrain/effects/health.html |date=January 18, 2008 }}. Epa.gov (June 2, 2006). Retrieved on February 9, 2013.</ref> Further, this form of pollution is linked to ] and ], with a study of 74,671 pregnant women in Beijing finding that every additional 100 μg/m<sup>3</sup> of {{SO2}} in the air reduced infants' weight by 7.3 g, making it and other forms of air pollution the largest attributable risk factor for low birth weight ever observed.<ref>{{Cite journal |last1=Wang |first1=X. |last2=Ding |first2=H. |last3=Ryan |first3=L. |last4=Xu |first4=X. |s2cid=2707126 |date=May 1, 1997 |title=Association between air pollution and low birth weight: a community-based study |journal=Environmental Health Perspectives |volume=105 |issue=5 |pages=514–20 |issn=0091-6765 |pmc=1469882 |pmid=9222137 |doi=10.1289/ehp.97105514}}</ref> |
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====Control measures==== |
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]s. While no ] may reach Maximum Feasible Reductions (MFRs), all assume steep declines from today's levels. By 2019, sulfate emission reductions were confirmed to proceed at a very fast rate.<ref name=XuRamanathanVictor>{{Cite journal|last1=Xu|first1=Yangyang|last2=Ramanathan|first2=Veerabhadran|last3=Victor|first3=David G.|date=December 5, 2018|title=Global warming will happen faster than we think|journal=Nature|language=en|volume=564|issue=7734|pages=30–32 |url=https://www.researchgate.net/publication/329411074 |doi=10.1038/d41586-018-07586-5|pmid=30518902|bibcode=2018Natur.564...30X|doi-access=free}}</ref>]] |
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Due largely to the US EPA's ], the U.S. has had a 33% decrease in emissions between 1983 and 2002 (see table). This improvement resulted in part from ], a technology that enables SO<sub>2</sub> to be chemically bound in ]s burning sulfur-containing coal or petroleum. |
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{| class="wikitable" |
|
|
|- |
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|
! Year |
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|
! SO<sub>2</sub> |
|
|
|- |
|
|
| 1970 |
|
|
| {{convert|31161000|ST|Mt|sigfig=3}} |
|
|
|- |
|
|
| 1980 |
|
|
| {{convert|25905000|ST|Mt|sigfig=3}} |
|
|
|- |
|
|
| 1990 |
|
|
|{{convert|23678000|ST|Mt|sigfig=3}} |
|
|
|- |
|
|
| 1996 |
|
|
|{{convert|18859000|ST|Mt|sigfig=3}} |
|
|
|- |
|
|
| 1997 |
|
|
|{{convert|19363000|ST|Mt|sigfig=3}} |
|
|
|- |
|
|
| 1998 |
|
|
|{{convert|19491000|ST|Mt|sigfig=3}} |
|
|
|- |
|
|
| 1999 |
|
|
|{{convert|18867000|ST|Mt|sigfig=3}} |
|
|
|} |
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|
|
|
|
In particular, ] reacts with sulfur dioxide to form ]: |
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: CaO + SO<sub>2</sub> → CaSO<sub>3</sub> |
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Aerobic oxidation of the CaSO<sub>3</sub> gives CaSO<sub>4</sub>, ]. Most gypsum sold in Europe comes from flue-gas desulfurization. |
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To control sulfur emissions, dozens of methods with relatively high efficiencies have been developed for fitting of coal-fired power plants.<ref>{{Cite journal|last1=Lin|first1=Cheng-Kuan|last2=Lin|first2=Ro-Ting|last3=Chen|first3=Pi-Cheng|last4=Wang|first4=Pu|last5=De Marcellis-Warin|first5=Nathalie|last6=Zigler|first6=Corwin|last7=Christiani|first7=David C.|date=February 8, 2018|title=A Global Perspective on Sulfur Oxide Controls in Coal-Fired Power Plants and Cardiovascular Disease|journal=Scientific Reports|language=en|volume=8|issue=1|pages=2611 |doi=10.1038/s41598-018-20404-2|pmid=29422539|issn=2045-2322|pmc=5805744|bibcode=2018NatSR...8.2611L}}</ref> Sulfur can be removed from coal during burning by using limestone as a bed material in ].<ref>{{cite book |last=Lindeburg |first=Michael R. |title=Mechanical Engineering Reference Manual for the PE Exam |location=Belmont, C.A. |publisher=Professional Publications, Inc |year=2006 |pages=27–3 |isbn=978-1-59126-049-3}}</ref> |
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Sulfur can also be removed from fuels before burning, preventing formation of SO<sub>2</sub> when the fuel is burnt. The ] is used in refineries to produce sulfur as a byproduct. The ] has also been used to remove sulfur from fuel. ] processes using iron oxides can also be used, for example, Lo-Cat<ref>. gtp-merichem.com</ref> or Sulferox.<ref>. (December 2002) Report by SFA Pacific, Inc. prepared for U.S. Department of Energy (PDF) Retrieved on October 31, 2011.</ref> |
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Fuel additives such as ] additives and magnesium carboxylate may be used in marine engines to lower the emission of sulfur dioxide gases into the atmosphere.<ref>May, Walter R. {{Webarchive|url=https://web.archive.org/web/20150402130001/http://www.fuelspec.com/library/Marine%20Emissions%20Abatement.pdf |date=April 2, 2015 }}. SFA International, Inc., p. 6.</ref> |
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===Impact on climate change=== |
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{{excerpt|Global dimming#History|paragraph=2}} |
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{{excerpt|Global dimming#Causes|paragraph=1|hat=no|files=no}} |
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====Projected impacts==== |
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], including the cooling provided by sulfate aerosols and the dimming they cause. The large ] shows that there are still substantial unresolved uncertainties.]] |
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{{excerpt|Global dimming#Future|paragraphs=1,3|hat=no|files=no}} |
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====Solar geoengineering==== |
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] into the stratosphere]] |
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{{excerpt|Stratospheric aerosol injection#Pollution controls and the discovery of radiative effects|paragraph=3|files=no}} |
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== Properties == |
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Table of thermal and physical properties of saturated liquid sulfur dioxide:<ref>{{Cite book |last=Holman |first=Jack P. |title=Heat Transfer |publisher=McGraw-Hill Companies, Inc. |year=2002 |isbn=9780072406559 |edition=9th |location=New York, NY |pages=600–606 |language=English}}</ref><ref>{{Cite book |last1=Incropera |last2=Dewitt |last3=Bergman |last4=Lavigne |first1=rank P. |first2=David P. |first3=Theodore L. |first4=Adrienne S. |title=Fundamentals of Heat and Mass Transfer |publisher=John Wiley and Sons, Inc. |year=2007 |isbn=9780471457282 |edition=6th |location=Hoboken, NJ |pages=941–950 |language=English}}</ref> |
|
|
{|class="wikitable mw-collapsible" |
|
|
|Temperature (°C) |
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|
|Density (kg/m^3) |
|
|
|Specific heat (kJ/kg K) |
|
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|Kinematic viscosity (m^2/s) |
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|Conductivity (W/m K) |
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|Thermal diffusivity (m^2/s) |
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|Prandtl Number |
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|Bulk modulus (K^-1) |
|
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|- |
|
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| −50 |
|
|
|1560.84 |
|
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|1.3595 |
|
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|4.84E-07 |
|
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|0.242 |
|
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|1.14E-07 |
|
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|4.24 |
|
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| – |
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|- |
|
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| −40 |
|
|
|1536.81 |
|
|
|1.3607 |
|
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|4.24E-07 |
|
|
|0.235 |
|
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|1.13E-07 |
|
|
|3.74 |
|
|
| – |
|
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|- |
|
|
| −30 |
|
|
|1520.64 |
|
|
|1.3616 |
|
|
|3.71E-07 |
|
|
|0.23 |
|
|
|1.12E-07 |
|
|
|3.31 |
|
|
| – |
|
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|- |
|
|
| −20 |
|
|
|1488.6 |
|
|
|1.3624 |
|
|
|3.24E-07 |
|
|
|0.225 |
|
|
|1.11E-07 |
|
|
|2.93 |
|
|
| – |
|
|
|- |
|
|
| −10 |
|
|
|1463.61 |
|
|
|1.3628 |
|
|
|2.88E-07 |
|
|
|0.218 |
|
|
|1.10E-07 |
|
|
|2.62 |
|
|
| – |
|
|
|- |
|
|
|0 |
|
|
|1438.46 |
|
|
|1.3636 |
|
|
|2.57E-07 |
|
|
|0.211 |
|
|
|1.08E-07 |
|
|
|2.38 |
|
|
| – |
|
|
|- |
|
|
|10 |
|
|
|1412.51 |
|
|
|1.3645 |
|
|
|2.32E-07 |
|
|
|0.204 |
|
|
|1.07E-07 |
|
|
|2.18 |
|
|
| – |
|
|
|- |
|
|
|20 |
|
|
|1386.4 |
|
|
|1.3653 |
|
|
|2.10E-07 |
|
|
|0.199 |
|
|
|1.05E-07 |
|
|
|2 |
|
|
|1.94E-03 |
|
|
|- |
|
|
|30 |
|
|
|1359.33 |
|
|
|1.3662 |
|
|
|1.90E-07 |
|
|
|0.192 |
|
|
|1.04E-07 |
|
|
|1.83 |
|
|
| – |
|
|
|- |
|
|
|40 |
|
|
|1329.22 |
|
|
|1.3674 |
|
|
|1.73E-07 |
|
|
|0.185 |
|
|
|1.02E-07 |
|
|
|1.7 |
|
|
| – |
|
|
|- |
|
|
|50 |
|
|
|1299.1 |
|
|
|1.3683 |
|
|
|1.62E-07 |
|
|
|0.177 |
|
|
|9.99E-08 |
|
|
|1.61 |
|
|
| – |
|
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|} |
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==See also== |
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* ] |
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* ] |
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* ] |
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* ] |
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==References== |
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{{Reflist|30em}} |
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==External links== |
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{{Commons category|Sulfur dioxide|lcfirst=yes}} |
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* |
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* |
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* |
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* |
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* |
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* |
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* |
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{{Oxides}} |
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{{Molecules detected in outer space}} |
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{{sulfur compounds}} |
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{{Authority control}} |
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] |
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] |
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] |
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