Sulfur dioxide
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Overview
Sulfur dioxide (also sulphur dioxide) is the chemical compound with the formula SO2. This important gas is the main product from the combustion of sulfur compounds and is of significant environmental concern. SO2 is often described as the "smell of burning sulfur" but is not responsible for the smell of rotten eggs.
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]
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 cement: CaSiO3 and 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.
Reactions
Treatment of basic solutions with sulfur dioxide affords sulfite salts:
- SO2 + 2 NaOH → Na2SO3 + H2O
Featuring sulfur in the +4 oxidation state, sulfur doxide 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
Uses
As a preservative
Sulfur dioxide is sometimes used as a preservative for dried apricots and other dried fruits due to its antimicrobial properties. The preservative is used to maintain the appearance of the fruit rather than to 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.[2] It is present even in so-called unsulphurated wine at concentrations of up to 10 milligrams per litre.[3] 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 doxide 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 evaporartion, 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.[4]
Emissions
According to the U.S. EPA (as presented by the 2002 World Almanac or in chart form[5]), 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[6].
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 solublity 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 solublity 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
According to Claude Ribbe in The Crime of Napoleon, sulfur dioxide gas was used as an execution poison by the French emperor to suppress a slave revolt in Haiti early in the 19th century.
References
- ↑ Holleman, A. F.; Wiberg, E. "Inorganic Chemistry" Academic Press: San Diego, 2001. ISBN 0-12-352651-5.
- ↑ 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.
- ↑ National Trends in Sulfur Dioxide Levels, United States Environmental Protection Agency.
- ↑ China has its worst spell of acid rain, United Press International.
See also
- Sulfur
- 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
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