Sulphur dioxide | |
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Other names | sulfur(IV) oxide; sulfurous anhydride |
Identifiers | |
CAS number | 7446-09-5 |
RTECS number | WS4550000 |
ChemSpider ID | |
Properties | |
Molecular formula | SO2 |
Molar mass | 64.07 g mol−1 |
Appearance | colourless gas |
Density | 2.551 g/L, gas 1.354 g/cm³ at -30°C 1.434 g/cm³ at 0°C 1.25 g/mL at 25°C |
Melting point |
−72.4 °C (200.75 K) |
Boiling point |
−10 °C (263 K) |
Solubility in water | 9.4 g/100 mL (25 °C) |
Vapor pressure | –10°C : 1.013 bar 20°C : 3.3 bar 40°C : 4.4 bar |
Acidity (pKa) | 1.81 |
Structure | |
Molecular shape | Bent 120°[1] |
Dipole moment | 1.63 D |
Hazards | |
EU classification | Toxic |
NFPA 704 |
3
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R-phrases | R23 R34 |
S-phrases | (S1/2) S9 S26 S36/37/39 S45 |
LC50 | 3000 ppm/30 min mouse inhalation |
Flash point | non-flammable |
Related compounds | |
Related compounds | Sulfur trioxide; sulfuric acid |
Except where noted otherwise, data are given for materials in their standard state (at 25 °C, 100 kPa) Infobox references |
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.[2] This is one of the causes for concern over the environmental impact of the use of these fuels as power sources.
Contents |
Sulfur dioxide can be prepared by burning sulfur:
The combustion of hydrogen sulfide and organosulfur compounds proceeds similarly.
The roasting of sulfide ores such as iron pyrites, sphalerite (zinc blende) and cinnabar (mercury sulfide) also releases SO2:
Sulfur dioxide is a by-product in the manufacture of calcium silicate cement: CaSO4 is heated with coke and sand in this process:
Action of hot sulfuric acid on copper turnings produces sulfur dioxide.
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[3]
Treatment of basic solutions with sulfur dioxide affords sulfite salts:
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:
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:
Sulfur dioxide can act as a metal binding ligand, typically where the transition metal is in oxidation state 0 or +1.[4] Up to 9 different bonding modes have been determined which include[4]:
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.
Sulfur dioxide is a very important compound in winemaking, and is designated as parts per million in wine, E number: E220.[5] It is present even in so-called unsulphurated 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. 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.
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.
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.
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.
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.
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.[7]
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. [8]
According to the U.S. EPA (as presented by the 2002 World Almanac or in chart form[9]), 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:
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[10].
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.
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) |
Sulfur dioxide acts as an acid, it reacts with other chemicals in the air to form tiny sulfate particles. When these are breathed, they gather in the lungs and are associated with increased respiratory symptoms and disease, difficulty in breathing, and premature death[11]. When mixed with water and contacted by skin, frostbite may occur. Upon contact with eyes, redness and pain will occur.[12].