|UN number||1079, 2037|
|Jmol-3D images||Image 1|
|Molar mass||64.066 g mol−1|
|Odor||Just-struck match/eggs like|
|Density||2.6288 kg m−3|
|Melting point||−72 °C; −98 °F; 201 K|
|Boiling point||−10 °C (14 °F; 263 K)|
|Solubility in water||94 g/L|
|Vapor pressure||237.2 kPa|
|Viscosity||0.403 cP (at 0 °C)|
|Dipole moment||1.62 D|
|248.223 J K−1 mol−1|
|Std enthalpy of
|-296.81 kJ mol−1|
|R-phrases||R23, R34, R50|
|S-phrases||(S1/2), S9, S26, S36/37/39, S45|
|LD50||3000 ppm (30 min inhaled, mouse)|
|Related sulfur oxides||Sulfur monoxide
|Except where noted otherwise, data are given for materials in their standard state (at 25 °C (77 °F), 100 kPa)|
|(what is: / ?)|
Sulfur dioxide (also sulphur dioxide) is the chemical compound with the formula SO
2. At standard atmosphere it is a toxic gas with a pungent, irritating and rotten smell. The triple point is 197.69 K and 1.67Kpa. It is released naturally by volcanic activity.
The specific person who discovered sulfur dioxide is not known, but it was used by the Romans in winemaking, when they discovered that burning sulfur candles inside empty wine vessels kept them fresh and free from vinegar smell.
- 1 Structure and bonding
- 2 Occurrence
- 3 Production
- 4 Reactions
- 5 Uses
- 6 As an air pollutant
- 7 Safety
- 8 See also
- 9 References
- 10 External links
Structure and bonding
SO2 is a bent molecule with C2v symmetry point group. A valence bond theory approach which just considered s and p orbitals, would describe the bonding in terms of resonance between two resonance structures.
The sulfur - oxygen bond has a bond order of 1.5. There is support for this simple approach that does not invoke d orbital participation. In terms of electron-counting formalism, the sulfur atom has an oxidation state of +4 and a formal charge of +1.
On other planets, it can be found in various concentrations, the most significant case being the atmosphere of Venus, where it is the third most significant atmospheric gas at 150 ppm. There it condenses to form clouds, and is a key component of chemical reactions in the planet's atmosphere and contributes to global warming. It has been implicated as a key agent in the warming of early Mars, with estimates of concentrations in the lower atmosphere as high as 100ppm though it currently only exists in trace amounts. On both Venus and Mars, its primary source, like on Earth, is believed to be volcanic. It is also believed to exist in trace amounts in the atmosphere of Jupiter and atmosphere of Saturn.
As an ice, it is thought to exist in abundance on the Galilean moons – as sublimating ice or frost on the trailing hemisphere of Io, a natural satellite of Jupiter and in the crust and mantle of Europa, Ganymede and Callisto possibly also in liquid form and readily reacting with water.
Sulfur dioxide is primarily produced for sulfuric acid manufacture (see contact process). In the United States in 1979, 23.6 million tonnes of sulfur dioxide was used in this way, compared with 150 thousand tonnes used for other purposes. Most sulfur dioxide is produced by the combustion of elemental sulfur. Some sulfur dioxide is also produced by roasting pyrite and other sulfide ores in air.
Sulfur dioxide is the product of the burning of sulfur or of burning materials that contain sulfur:
- S + O2 → SO2, ΔH = -297 kJ/mol
To aid combustion, liquefied sulfur (140–150 °C) is sprayed through an atomizing nozzle to generate fine drops of sulfur with a large surface area. The reaction is exothermic, and the combustion produces temperatures of 1000–1600 °C. The significant amount of heat produced is recovered by steam generation that can subsequently be converted to electricity.
The combustion of hydrogen sulfide and organosulfur compounds proceeds similarly. For example:
- 2 H2S + 3 O2 → 2 H2O + 2 SO2
- 4 FeS2 + 11 O2 → 2 Fe2O3 + 8 SO2
- 2 ZnS + 3 O2 → 2 ZnO + 2 SO2
- HgS + O2 → Hg + SO2
- 4 FeS + 7O2 → 2 Fe2O3 + 4 SO2
A combination of these reactions is responsible for the largest source of sulfur dioxide, volcanic eruptions. These events can release millions of tonnes of SO2.
Reduction of higher oxides
- 2 CaSO4 + 2 SiO2 + C → 2 CaSiO3 + 2 SO2 + CO2
Up until the 1970s, commercial quantities of sulfuric acid and cement were produced by this process in Whitehaven, England. Upon being mixed with shale or marl, and roasted, the sulfate liberated sulfur dioxide gas, used in sulfuric acid production, the reaction also produced calcium silicate, a precursor in cement production.
- Cu + 2 H2SO4 → CuSO4 + SO2 + 2 H2O
Sulfite results from the reaction of aqueous base and sulfur dioxide. The reverse reaction involves acidification of sodium metabisulfite:
- H2SO4 + Na2S2O5 → 2 SO2 + Na2SO4 + H2O
Treatment of basic solutions with sulfur dioxide affords sulfite salts:
- SO2 + 2 NaOH → Na2SO3 + H2O
- SO2 + Cl2 → SO2Cl2
- SO2 + 2 H2S → 3 S + 2 H2O
The sequential oxidation of sulfur dioxide followed by its hydration is used in the production of sulfuric acid.
- 2 SO2 + 2 H2O + O2 → 2 H2SO4
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 dichromate solution, turning the solution from orange to green (Cr3+ (aq)). It can also reduce ferric ions to ferrous
Sulfur dioxide can bind to metal ions as a ligand to form metal sulfur dioxide complexes, 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 η1.
Precursor to sulfuric acid
Sulfur dioxide is an intermediate in the production of sulfuric acid, being converted to sulfur trioxide, and then to oleum, which is made into sulfuric acid. Sulfur dioxide for this purpose is made when sulfur combines with oxygen. The method of converting sulfur dioxide to sulfuric acid is called the contact process. Several billion kilograms are produced annually for this purpose.
As a preservative
Sulfur dioxide is sometimes used as a preservative for dried apricots, dried figs, and other dried fruits owing to its antimicrobial properties, and it is sometimes called E220 when used in this way. As a preservative, it maintains the colorful appearance of the fruit and prevents rotting. It is also added to sulfured molasses.
The specific person who discovered sulfur dioxide is not known, but it was first used by the Romans in winemaking, when they discovered that burning sulfur candles inside empty wine vessels keeps them fresh and free from vinegar smell.
Sulfur dioxide is an important compound in winemaking, and is designated as parts per million in wine, E number: E220. It is present even in so-called unsulfurated wine at concentrations of up to 10 mg/L. It serves as an antibiotic and antioxidant, protecting wine from spoilage by bacteria and oxidation. Its antimicrobial action also helps to minimize volatile acidity. Sulfur dioxide is responsible for the words "contains sulfites" found on wine labels.
Sulfur dioxide exists in wine in free and bound forms, and the combinations are referred to as total SO2. Binding, for instance to the carbonyl group of acetaldehyde, varies with the wine in question. The free form exists in equilibrium between molecular SO2 (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) SO2, which is the active form, while at higher pH more SO2 is found in the inactive sulfite and bisulfite forms. It is the molecular SO2 which is active as an antimicrobial and antioxidant, and this is also the form which may be perceived as a pungent odour at high levels. Wines with total SO2 concentrations below 10 ppm do not require "contains sulfites" on the label by US and EU laws. The upper limit of total SO2 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, SO2 is mostly undetectable in wine, but at free SO2 concentrations over 50 ppm, SO2 becomes evident in the nose and taste of wine.
SO2 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 the risk of cork taint, a mixture of SO2, water, and citric acid is commonly used to clean and sanitize equipment. Compounds of ozone (O3) are now used extensively as cleaning products in wineries due to their efficiency, and because these compounds do not affect the wine or equipment.
As a reducing agent
Sulfur dioxide is also a good reductant. In the presence of water, sulfur dioxide is able to decolorize substances. Specifically it is a useful reducing bleach for papers and delicate materials such as clothes. This bleaching effect normally does not last very long. Oxygen in the atmosphere reoxidizes the reduced dyes, restoring the color. In municipal wastewater treatment, sulfur dioxide is used to treat chlorinated wastewater prior to release. Sulfur dioxide reduces free and combined chlorine to chloride.
Sulfur dioxide is fairly soluble in water, and by both IR and Raman spectroscopy, it is known that the hypothetical sulfurous acid, H2SO3, is not present to any extent. However, such solutions do show spectra of the hydrogen sulfite ion, HSO3−, by reaction with water, and it is in fact the actual reducing agent present:
- SO2 + H2O ⇌ HSO3− + H+
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. The role of sulfur dioxide in mammalian biology is not yet well understood. Sulfur dioxide blocks nerve signals from the pulmonary stretch receptors (PSRs) and abolishes the Hering–Breuer inflation reflex.
As a refrigerant
Being easily condensed and possessing a high heat of evaporation, sulfur dioxide is a candidate material for refrigerants. Prior to the development of CFCs, sulfur dioxide was used as a refrigerant in home refrigerators.
As a reagent and solvent in the laboratory
Sulfur dioxide is a versatile inert solvent that has been widely used for dissolving highly oxidizing salts. It is also used occasionally as a source of the sulfonyl group in organic synthesis. Treatment of aryl diazonium salts with sulfur dioxide and cuprous chloride affords the corresponding aryl sulfonyl chloride, for example:
As an air pollutant
Sulfur dioxide is a noticeable component in the atmosphere, especially following volcanic eruptions. According to the United States Environmental Protection Agency (as presented by the 2002 World Almanac or in chart form), the following amount of sulfur dioxide was released in the U.S. per year:
|1970||31,161,000 short tons (28.3 Mt)|
|1980||25,905,000 short tons (23.5 Mt)|
|1990||23,678,000 short tons (21.5 Mt)|
|1996||18,859,000 short tons (17.1 Mt)|
|1997||19,363,000 short tons (17.6 Mt)|
|1998||19,491,000 short tons (17.7 Mt)|
|1999||18,867,000 short tons (17.1 Mt)|
Sulfur dioxide is a major air pollutant and has significant impacts upon human health. In addition the concentration of sulfur dioxide in the atmosphere can influence the habitat suitability for plant communities as well as animal life. Sulfur dioxide emissions are a precursor to acid rain and atmospheric particulates. Due largely to the US EPA’s Acid Rain Program, the U.S. has witnessed a 33% decrease in emissions between 1983 and 2002. This improvement resulted in part from flue-gas desulfurization, a technology that enables SO2 to be chemically bound in power plants burning sulfur-containing coal or oil. In particular, calcium oxide (lime) reacts with sulfur dioxide to form calcium sulfite:
- CaO + SO2 → CaSO3
Aerobic oxidation of the CaSO3 gives CaSO4, anhydrite. Most gypsum sold in Europe comes from flue-gas desulfurization.
Sulfur can also be removed from fuels prior to burning the fuel, preventing the formation of SO2 because there is no sulfur in the fuel from which SO2 can be formed. The Claus process is used in refineries to produce sulfur as a byproduct. The Stretford process has also been used to remove sulfur from fuel. Redox processes using iron oxides can also be used, for example, Lo-Cat or Sulferox.
As of 2006, China was the world's largest sulfur dioxide polluter, with 2005 emissions estimated to be 25,490,000 short tons (23.1 Mt). This amount represents a 27% increase since 2000, and is roughly comparable with U.S. emissions in 1980.
Inhaling sulfur dioxide is associated with increased respiratory symptoms and disease, difficulty in breathing, and premature death. In 2008, the American Conference of Governmental Industrial Hygienists reduced the short-term exposure limit to 5 parts per million (ppm). The OSHA PEL is currently set at 5 ppm (13 mg/m3) time weighted average. NIOSH has set the IDLH at 100 ppm. In 2010, the EPA "revised the primary SO2 NAAQS by establishing a new 1-hour standard at a level of 75 parts per billion (ppb). EPA revoked the two existing primary standards because they would not provide additional public health protection given a 1-hour standard at 75 ppb."
In the United States, the Center for Science in the Public Interest lists the two food preservatives, sulfur dioxide and sodium bisulfite, as being safe for human consumption except for certain asthmatic individuals who may be sensitive to them, especially in large amounts. Symptoms of sensitivity to sulfiting agents, including sulfur dioxide, manifest as potentially life-threatening trouble breathing within minutes of ingestion.
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|Wikimedia Commons has media related to sulfur dioxide.|
- United States Environmental Protection Agency Sulfur Dioxide page
- International Chemical Safety Card 0074
- IARC Monographs. "Sulfur Dioxide and some Sulfites, Bisulfites and Metabisulfites" v54. 1992. p131.
- NIOSH Pocket Guide to Chemical Hazards
- CDC – Sulfure Dioxide – NIOSH Workplace Safety and Health Topic
- Sulfur Dioxide, Molecule of the Month