Sulfur dioxide: Difference between revisions
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'''Sulfur dioxide''' (also '''sulphur dioxide''') is the [[chemical compound]] with the formula SO<sub>2</sub>. SO<sub>2</sub> is produced by [[volcano]]es and in various industrial processes. Since [[coal]] and [[petroleum]] often contain sulfur compounds, their combustion generates sulfur dioxide. Further oxidation of SO<sub>2</sub>, usually in the presence of a catalyst such as [[nitrogen dioxide|NO<sub>2</sub>]], forms [[sulfuric acid|H<sub>2</sub>SO<sub>4</sub>]], and thus [[acid rain]].<ref>Holleman, A. F.; Wiberg, E. "Inorganic Chemistry" Academic Press: San Diego, 2001. ISBN 0-12-352651-5.</ref> This is one of the causes for concern over the environmental impact of the use of these fuels as power sources. |
'''Sulfur dioxide''' (also '''sulphur dioxide''') is the [[chemical compound]] with the formula SO<sub>2</sub>. SO<sub>2</sub> is produced by [[volcano]]es and in various industrial processes. Since [[coal]] and [[petroleum]] often contain sulfur compounds, their combustion generates sulfur dioxide. Further oxidation of SO<sub>2</sub>, usually in the presence of a catalyst such as [[nitrogen dioxide|NO<sub>2</sub>]], forms [[sulfuric acid|H<sub>2</sub>SO<sub>4</sub>]], and thus [[acid rain]].<ref>Holleman, A. F.; Wiberg, E. "Inorganic Chemistry" Academic Press: San Diego, 2001. ISBN 0-12-352651-5.</ref> This is one of the causes for concern over the environmental impact of the use of these fuels as power sources. |
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==Preparation== |
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smelly fanny fingers are good in biscuit bread ummmm |
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Sulfur dioxide can be prepared by burning [[sulfur]]: |
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oh yeah acids and the answer is suluric acid makes acid rain!!! have fun xxx |
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:S<sub>8</sub> + 8 O<sub>2</sub> → 8 SO<sub>2</sub> |
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The combustion of [[hydrogen sulfide]] and organosulfur compounds proceeds similarly. |
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: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) |
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The roasting of sulfide ores such as iron [[pyrites]], [[sphalerite]] (zinc blende) and [[cinnabar]] (mercury sulfide) also releases SO<sub>2</sub>: |
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:4 [[Iron|Fe]]S<sub>2</sub>(s) + 11 O<sub>2</sub>(g) → 2 Fe<sub>2</sub>O<sub>3</sub>(s) + 8 SO<sub>2</sub>(g) |
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:2 [[Zinc|Zn]]S(s) + 3 O<sub>2</sub>(g) → 2 ZnO(s) + 2 SO<sub>2</sub>(g) |
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:HgS(s) + O<sub>2</sub>(g) → Hg(g) + SO<sub>2</sub>(g) |
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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: |
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:2 CaSO<sub>4</sub>(s) + 2SiO<sub>2</sub>(s) + C(s) → 2 CaSiO<sub>3</sub>(s) + 2 SO<sub>2</sub>(g) + CO<sub>2</sub>(g) |
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Action of hot [[sulfuric acid]] on copper [[swarf|turnings]] produces sulfur dioxide. |
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:Cu(s) + 2H<sub>2</sub>SO<sub>4</sub>(aq) → CuSO<sub>4</sub>(aq) + SO<sub>2</sub>(g) + 2H<sub>2</sub>O(l) |
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==Structure and bonding== |
==Structure and bonding== |
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Revision as of 11:26, 10 October 2008
Template:Chembox new Sulfur dioxide (also sulphur dioxide) is the chemical compound with the formula SO2. SO2 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.[1] 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 + 8 O2 → 8 SO2
The combustion of hydrogen sulfide and organosulfur compounds proceeds similarly.
- 2 H2S(g) + 3 O2(g) → 2 H2O(g) + 2 SO2(g)
The roasting of sulfide ores such as iron pyrites, sphalerite (zinc blende) and cinnabar (mercury sulfide) also releases SO2:
- 4 FeS2(s) + 11 O2(g) → 2 Fe2O3(s) + 8 SO2(g)
- 2 ZnS(s) + 3 O2(g) → 2 ZnO(s) + 2 SO2(g)
- HgS(s) + O2(g) → Hg(g) + SO2(g)
Sulfur dioxide is a by-product in the manufacture of calcium silicate cement: CaSO4 is heated with coke and sand in this process:
- 2 CaSO4(s) + 2SiO2(s) + C(s) → 2 CaSiO3(s) + 2 SO2(g) + CO2(g)
Action of hot sulfuric acid on copper turnings produces sulfur dioxide.
- Cu(s) + 2H2SO4(aq) → CuSO4(aq) + SO2(g) + 2H2O(l)
Structure and bonding
SO2 is a bent molecule with C2v symmetry point group. 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.
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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 O3 (127.8 pm) than in dioxygen, O2 (120.7 pm). The mean bond energy is greater in SO2 (548 kJ mol−1) than in SO (524 kJ mol−1), whereas it is less in O3 (297 kJ mol−1) than in O2 (490 kJ mol−1). 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[2]
Reactions
Treatment of basic solutions with sulfur dioxide affords sulfite salts:
- SO2 + 2 NaOH → Na2SO3 + H2O
Featuring sulfur in the +4 oxidation state, sulfur dioxide is a reducing agent. It is oxidized by halogens such as chlorine to give the sulfuryl halides:
- 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 + 2 H2S → 3 S + 2 H2O
Sulfur dioxide can act as a metal binding ligand, typically where the transition metal is in oxidation state 0 or +1.[3] Up to 9 different bonding modes have been determined which include[3]:
- S donation -planar and pyramidal η1
- O donation η1
- π donation- side on η2
- S bridging across two metal centres or two ends of a metal-metal bon
- O-S bridging to two metal centres
- bridging- one metal centre π donation- side on, the other metal center O donation
- bridging over three metal centres
Uses
As a preservative
Sulfur dioxide is sometimes used as a preservative for dried apricots and other dried fruits due to its antimicrobial properties, it is sometimes called E220 when used in this way. The preservative is used to maintain the appearance of the fruit and prevent rotting. Its presence can give fruit a distinctive chemical taste.
In winemaking
Sulfur dioxide is a very important compound in winemaking, and is designated as parts per million in wine, E number: E220.[4] It is present even in so-called unsulphurated wine at concentrations of up to 10 milligrams per litre.[5] 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 SO2 concentrations below 10ppm do not require "contains sulfites" on the label by US and EU laws. The upper limit of SO2 allowed in wine is 350ppm in US, in the EU 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 very clean, and because bleach cannot be used in a winery, a mixture of SO2, water, and citric acid is commonly used to clean hoses, tanks, and other equipment to keep it clean and free of bacteria.
As a reducing bleach
Sulfur dioxide is also a good 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.
Precursor to sulfuric acid
Sulfur dioxide is also used to make sulfuric acid, being converted to sulfur trioxide, and then to oleum, which is made into sulfuric acid. Sulfur dioxide for this purpose is made when sulfur combines with oxygen. The method of converting sulfur dioxide to sulfuric acid is called the contact process.
Biochemical and biomedical roles
Sulfur dioxide is toxic in large amounts. It or its conjugate base bisulfite is produced biologically as an intermediate in both sulfate-reducing organisms and in sulfur oxidizing bacteria as well. Sulfur dioxide has no role in mammalian biology. Sulfur dioxide blocks nerve signals from the pulmonary stretch receptors (PSR's) and abolishes the Hering-Breuer inflation reflex.
As a refrigerant
Being easily condensed and with a high heat of evaporation, sulfur dioxide is a candidate material for refrigerants. Prior to the development of freons, sulfur dioxide was used as a refrigerant in home refrigerators.
As a reagent and solvent
Sulfur dioxide is a versatile inert solvent that has been widely used for dissolving highly oxidizing salts. It is also used occasionally as a source of the sulfonyl group in organic synthesis. Treatment of aryldiazonium salts with sulfur dioxide affords the corresponding aryl sulfonyl chloride.[6]
Dechlorination
In municipal wastewater treatment sulfur dioxide is used to treat chlorinated wastewater prior to release. Sulfur dioxide reacts with free and combined chlorine to form negatively charged chloride ions. [7]
Emissions
According to the U.S. EPA (as presented by the 2002 World Almanac or in chart form[8]), the following amount of sulfur dioxide was released in the U.S. per year, measured in thousands of short tons:
| *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 converts this CaSO3 into 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[9].
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
| 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.
Threats to Health
Sulfur dioxide acts as an acid. Inhalation results in labored breathing, coughing, and/or a sore throat and may cause permanent pulmonary damage. When mixed with water and contacted by skin, frostbite may occur. When it makes contact with eyes, redness and pain will occur.[10]
References
- ^ Holleman, A. F.; Wiberg, E. "Inorganic Chemistry" Academic Press: San Diego, 2001. ISBN 0-12-352651-5.
- ^ Greenwood, Norman N.; Earnshaw, Alan (1997). Chemistry of the Elements (2nd ed.). Butterworth-Heinemann. ISBN 978-0-08-037941-8. p. 700
- ^ a b Greenwood, Norman N.; Earnshaw, Alan (1997). Chemistry of the Elements (2nd ed.). Butterworth-Heinemann. ISBN 978-0-08-037941-8.
- ^ Current EU approved additives and their E Numbers, The Food Standards Agency website.
- ^ Sulphites in wine, MoreThanOrganic.com.
- ^ R. V. Hoffman “m-Trifluoromethylbenzenesulfonyl chloride” Organic Syntheses, Collected Volume 7, p.508 (1990). http://www.orgsyn.org/orgsyn/pdfs/CV7P0508.pdf.
- ^ Tchobanoglous, George. Wastewater Engineering. 3rd ed. New York: Mc Graw Hill, 1979.
- ^ National Trends in Sulfur Dioxide Levels, United States Environmental Protection Agency.
- ^ China has its worst spell of acid rain, United Press International.
- ^ SULPHUR DIOXIDE, International Labour Organization.
See also
- Sulfur trioxide
- Sulfur-iodine cycle
- National Ambient Air Quality Standards
- Homer City Generating Station
External links
- United States Environmental Protection Agency Sulfur Dioxide page
- International Chemical Safety Card 0074
- IARC Monograph "Sulfur Dioxide and some Sulfites, Bisulfites and Metabisulfites"
- NIOSH Pocket Guide to Chemical Hazards
- Food Intolerance Network - Sulfite factsheet
- Sulfur Dioxide, Molecule of the Month