Sulfur dioxide (also sulphur dioxide in British English) is the chemical compound with the formula SO
2. It is a toxic gas with a burnt match smell. It is released naturally by volcanic activity and is produced as a by-product of the burning of fossil fuels contaminated with sulfur compounds and copper extraction.
3D model (JSmol)
|E number||E220 (preservatives)|
|UN number||1079, 2037|
|Molar mass||64.066 g mol−1|
|Odor||Pungent; similar to a just-struck match|
|Density||2.6288 kg m−3|
|Melting point||−72 °C; −98 °F; 201 K|
|Boiling point||−10 °C (14 °F; 263 K)|
forms sulfurous acid
|Vapor pressure||237.2 kPa|
|Viscosity||0.403 cP (at 0 °C)|
|248.223 J K−1 mol−1|
Std enthalpy of
|−296.81 kJ mol−1|
|GHS signal word||Danger|
|Lethal dose or concentration (LD, LC):|
LC50 (median concentration)
|3000 ppm (mouse, 30 min)|
2520 ppm (rat, 1 hr)
LCLo (lowest published)
|993 ppm (rat, 20 min)|
611 ppm (rat, 5 hr)
764 ppm (mouse, 20 min)
1000 ppm (human, 10 min)
3000 ppm (human, 5 min)
|US health exposure limits (NIOSH):|
|TWA 5 ppm (13 mg/m3)|
|TWA 2 ppm (5 mg/m3) ST 5 ppm (13 mg/m3)|
IDLH (Immediate danger)
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).
SO2 is a bent molecule with C2v symmetry point group. A valence bond theory approach considering just 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 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 100 ppm, though it only exists in trace amounts. On both Venus and Mars, as on Earth, its primary source is thought to be volcanic. The atmosphere of Io, a natural satellite of Jupiter, is 90% sulfur dioxide and trace amounts are thought to also exist in the atmosphere of Jupiter.
As an ice, it is thought to exist in abundance on the Galilean moons—as subliming ice or frost on the trailing hemisphere of Io, 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 (26,014,547 US short tons) of sulfur dioxide were used in this way, compared with 150 thousand tonnes (165,347 US short tons) 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:
To aid combustion, liquefied sulfur (140–150 °C, 284-302 °F) 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, (1832-2912 °F). 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:
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.
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.
On a laboratory scale, the action of hot concentrated sulfuric acid on copper turnings produces sulfur dioxide.
Sulfite results from the reaction of aqueous base and sulfur dioxide. The reverse reaction involves acidification of sodium metabisulfite:
Treatment of basic solutions with sulfur dioxide affords sulfite salts (e.g. sodium sulfite):
The sequential oxidation of sulfur dioxide followed by its hydration is used in the production of sulfuric acid.
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 react with certain 1,3-dienes in a cheletropic reaction to form cyclic sulfones. This reaction is exploited on an industrial scale for the synthesis of sulfolane, which is an important solvent in the petrochemical industry.
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.
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.
Sulfur dioxide is sometimes used as a preservative for dried apricots, dried figs, and other dried fruits, owing to its antimicrobial properties, and is called E220 when used in this way in Europe. As a preservative, it maintains the colorful appearance of the fruit and prevents rotting. It is also added to sulfured molasses.
It is still an important compound in winemaking, and is measured in parts per million (ppm) in wine. 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 - a phenomenon that leads to the browning of the wine and a loss of cultivar specific flavors. Its antimicrobial action also helps minimize volatile acidity. Wines containing sulfur dioxide are typically labeled with "containing sulfites".
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. The molecular SO2 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 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 smell 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. Ozone (O3) is now used extensively for sanitizing in wineries due to its efficacy, and because it does not affect the wine or most equipment.
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; 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:
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 and abolishes the Hering–Breuer inflation reflex.
It was shown that endogenous sulfur dioxide plays a role in diminishing an experimental lung damage caused by oleic acid. Endogenous sulfur dioxide lowered lipid peroxidation, free radical formation, oxidative stress and inflammation during an experimental lung damage. Conversely, a successful lung damage caused a significant lowering of endogenous sulfur dioxide production, and an increase in lipid peroxidation, free radical formation, oxidative stress and inflammation. Moreover, blockade of an enzyme that produces endogenous SO2 significantly increased the amount of lung tissue damage in the experiment. Conversely, adding acetylcysteine or glutathione to the rat diet increased the amount of endogenous SO2 produced and decreased the lung damage, the free radical formation, oxidative stress, inflammation and apoptosis.
It is considered that endogenous sulfur dioxide plays a significant physiological role in regulating cardiac and blood vessel function, and aberrant or deficient sulfur dioxide metabolism can contribute to several different cardiovascular diseases, such as arterial hypertension, atherosclerosis, pulmonary arterial hypertension, stenocardia.
It was shown that in children with pulmonary arterial hypertension due to congenital heart diseases the level of homocysteine 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.
Endogenous sulfur dioxide also has been shown to lower the proliferation rate of endothelial smooth muscle cells in blood vessels, via lowering the MAPK activity and activating adenylyl cyclase and protein kinase A. Smooth muscle cell proliferation is one of important mechanisms of hypertensive remodeling of blood vessels and their stenosis, so it is an important pathogenetic mechanism in arterial hypertension and atherosclerosis.
Endogenous sulfur dioxide in low concentrations causes endothelium-dependent vasodilation. 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 calcium channels 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.
Being easily condensed and possessing a high heat of evaporation, sulfur dioxide is a candidate material for refrigerants. Prior to the development of chlorofluorocarbons, sulfur dioxide was used as a refrigerant in home refrigerators.
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 organic synthesis. Treatment of aryl diazonium salts with sulfur dioxide and cuprous chloride yields the corresponding aryl sulfonyl chloride, for example:
Injections of sulfur dioxide in the stratosphere has been proposed in climate engineering. The cooling effect would be similar to what has been observed after the large explosive volcano eruption of Mount Pinatubo in 1991. However this form of geoengineering would have uncertain regional consequences on rainfall patterns, for example in monsoon regions.
Sulfur dioxide is a noticeable component in the atmosphere, especially following volcanic eruptions. According to the United States Environmental Protection Agency, the amount of sulfur dioxide released in the U.S. per year was:
|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 had 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:
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 before burning, preventing formation of SO2 when the fuel is burnt. 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 0.25 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 one-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 one-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.
The Acid Rain Program is a market-based initiative taken by the United States Environmental Protection Agency in an effort to reduce overall atmospheric levels of sulfur dioxide and nitrogen oxides, which cause acid rain. The program is an implementation of emissions trading that primarily targets coal-burning power plants, allowing them to buy and sell emission permits (called "allowances") according to individual needs and costs. In 2011, the trading program that existed since 1995 was supplemented by four separate trading programs under the Cross-State Air Pollution Rule (CSAPR). On August 21, 2012, the United States Court of Appeals for the District of Columbia issued its Opinion and Order in the appeal of the Cross State Air Pollution Rule (CSAPR) for two independent legal reasons. The stay on CSAPR was lifted in October 2014, allowing implementation of the law and its trading programs to begin.A 2017 NBER paper found that the "Acid Rain Program caused lasting improvements in ambient air quality," reducing mortality risk by 5% over 10 years.Acid rain
Acid rain is a rain or any other form of precipitation that is unusually acidic, meaning that it has elevated levels of hydrogen ions (low pH). It can have harmful effects on plants, aquatic animals and infrastructure. Acid rain is caused by emissions of sulfur dioxide and nitrogen oxide, which react with the water molecules in the atmosphere to produce acids. Some governments have made efforts since the 1970s to reduce the release of sulfur dioxide and nitrogen oxide into the atmosphere with positive results. Nitrogen oxides can also be produced naturally by lightning strikes, and sulfur dioxide is produced by volcanic eruptions. Acid rain has been shown to have adverse impacts on forests, freshwaters and soils, killing insect and aquatic life-forms, causing paint to peel, corrosion of steel structures such as bridges, and weathering of stone buildings and statues as well as having impacts on human health.Calcium bisulfite
Calcium bisulfite (calcium bisulphite) is an inorganic compound which is the salt of a calcium cation and a bisulfite anion. It may be prepared by reacting lime with an excess of sulfurous acid, essentially a mixture of sulfur dioxide and water. It is a weak reducing agent, as is sulfur dioxide, sulfites, and any other compound containing sulfur in the +4 oxidation state. As a food additive it is used as a preservative under the E number E227. Calcium bisulfite is an acid salt and behaves like an acid in aqueous solution.Contact process
The contact process is the current method of producing sulfuric acid in the high concentrations needed for industrial processes. Platinum used to be the catalyst for this reaction; however, as it is susceptible to reacting with arsenic impurities in the sulfur feedstock, vanadium(V) oxide (V2O5) is now preferred.This process was patented in 1831 by British vinegar merchant Peregrine Phillips. In addition to being a far more economical process for producing concentrated sulfuric acid than the previous lead chamber process, the contact process also produces sulfur trioxide and oleum.Flue-gas desulfurization
Flue-gas desulfurization (FGD) is a set of technologies used to remove sulfur dioxide (SO2) from exhaust flue gases of fossil-fuel power plants, and from the emissions of other sulfur oxide emitting processes (e.g trash incineration)Flue gas
Flue gas is the gas exiting to the atmosphere via a flue, which is a pipe or channel for conveying exhaust gases from a fireplace, oven, furnace, boiler or steam generator. Quite often, the flue gas refers to the combustion exhaust gas produced at power plants. Its composition depends on what is being burned, but it will usually consist of mostly nitrogen (typically more than two-thirds) derived from the combustion of air, carbon dioxide (CO2), and water vapor as well as excess oxygen (also derived from the combustion air). It further contains a small percentage of a number of pollutants, such as particulate matter (like soot), carbon monoxide, nitrogen oxides, and sulfur oxides.Most fossil fuels are combusted with ambient air (as differentiated from combustion with pure oxygen). Since ambient air contains about 79 volume percent gaseous nitrogen (N2), which is essentially non-combustible, the largest part of the flue gas from most fossil-fuel combustion is uncombusted nitrogen. Carbon dioxide (CO2), the next largest part of flue gas, can be as much as 10−25 volume percent or more of the flue gas. This is closely followed in volume by water vapor (H2O) created by the combustion of the hydrogen in the fuel with atmospheric oxygen. Much of the 'smoke' seen pouring from flue gas stacks is this water vapor forming a cloud as it contacts cool air.
A typical flue gas from the combustion of fossil fuels contains very small amounts of nitrogen oxides (NOx), sulfur dioxide (SO2) and particulate matter. The nitrogen oxides are derived from the nitrogen in the ambient air as well as from any nitrogen-containing compounds in the fossil fuel. The sulfur dioxide is derived from any sulfur-containing compounds in the fuels. The particulate matter is composed of very small particles of solid materials and very small liquid droplets which give flue gases their smoky appearance.
The steam generators in large power plants and the process furnaces in large refineries, petrochemical and chemical plants, and incinerators burn considerable amounts of fossil fuels and therefore emit large amounts of flue gas to the ambient atmosphere. The table below presents the total amounts of flue gas typically generated by the burning of fossil fuels such as natural gas, fuel oil and coal. The data were obtained by stoichiometric calculations.The total amount of wet flue gas generated by coal combustion is only 10 percent higher than the flue gas generated by natural-gas combustion (the ratio for dry flue gas is higher).Io (moon)
Io (Jupiter I) is the innermost of the four Galilean moons of the planet Jupiter. It is the fourth-largest moon, has the highest density of all the moons, and has the least amount of water of any known astronomical object in the Solar System. It was discovered in 1610 and was named after the mythological character Io, a priestess of Hera who became one of Zeus' lovers.
With over 400 active volcanoes, Io is the most geologically active object in the Solar System. This extreme geologic activity is the result of tidal heating from friction generated within Io's interior as it is pulled between Jupiter and the other Galilean satellites—Europa, Ganymede and Callisto. Several volcanoes produce plumes of sulfur and sulfur dioxide that climb as high as 500 km (300 mi) above the surface. Io's surface is also dotted with more than 100 mountains that have been uplifted by extensive compression at the base of Io's silicate crust. Some of these peaks are taller than Mount Everest. Unlike most satellites in the outer Solar System, which are mostly composed of water ice, Io is primarily composed of silicate rock surrounding a molten iron or iron-sulfide core. Most of Io's surface is composed of extensive plains coated with sulfur and sulfur-dioxide frost.
Io's volcanism is responsible for many of its unique features. Its volcanic plumes and lava flows produce large surface changes and paint the surface in various subtle shades of yellow, red, white, black, and green, largely due to allotropes and compounds of sulfur. Numerous extensive lava flows, several more than 500 km (300 mi) in length, also mark the surface. The materials produced by this volcanism make up Io's thin, patchy atmosphere and Jupiter's extensive magnetosphere. Io's volcanic ejecta also produce a large plasma torus around Jupiter.
Io played a significant role in the development of astronomy in the 17th and 18th centuries. It was discovered in January 1610 by Galileo Galilei, along with the other Galilean satellites. This discovery furthered the adoption of the Copernican model of the Solar System, the development of Kepler's laws of motion, and the first measurement of the speed of light. From Earth, Io remained just a point of light until the late 19th and early 20th centuries, when it became possible to resolve its large-scale surface features, such as the dark red polar and bright equatorial regions. In 1979, the two Voyager spacecraft revealed Io to be a geologically active world, with numerous volcanic features, large mountains, and a young surface with no obvious impact craters. The Galileo spacecraft performed several close flybys in the 1990s and early 2000s, obtaining data about Io's interior structure and surface composition. These spacecraft also revealed the relationship between Io and Jupiter's magnetosphere and the existence of a belt of high-energy radiation centered on Io's orbit. Io receives about 3,600 rem (36 Sv) of ionizing radiation per day.Further observations have been made by Cassini–Huygens in 2000, New Horizons in 2007, and Juno in 2017 and 2018, as well as from Earth-based telescopes and the Hubble Space Telescope.Metal sulfur dioxide complex
Metal sulfur dioxide complexes are complexes that contain sulfur dioxide, SO2, bonded to a transition metal. Such compounds are common but are mainly of theoretical interest. Historically, the study of these compounds has provided insights into the mechanisms of migratory insertion reactions in organometallic chemistry.Sodium bisulfite
Sodium bisulfite (or sodium bisulphite, sodium hydrogen sulfite) is a chemical compound with the chemical formula NaHSO3. Sodium bisulfite is a food additive with E number E222. This salt of bisulfite can be prepared by bubbling sulfur dioxide in a solution of sodium carbonate in water. Sodium bisulfite in contact with chlorine bleach (aqueous solution of sodium hypochlorite) will generate heat and form sodium bisulfate and sodium chloride.Sodium sulfite
Sodium sulfite (sodium sulphite) is a soluble sodium salt of sulfurous acid (sulfite) with the chemical formula Na2SO3. It is also used as a preservative to prevent dried fruit from discoloring, and for preserving meats, and is used in the same way as sodium thiosulfate to convert elemental halogens to their respective hydrohalic acids, in photography and for reducing chlorine levels in pools. In boiler systems, sulfite and bisulfite are the most commonly employed oxygen scavengers used to prevent pitting corrosion. Sodium sulfite is also a byproduct of sulfur dioxide scrubbing, a part of the flue-gas desulfurization process.Sulfinic acid
Sulfinic acids are oxoacids of sulfur with the structure RSO(OH). In these organosulfur compounds, sulfur is pyramidal.
They are often prepared in situ by acidification of the corresponding sulfinate salts, which are typically more robust than the acid. These salts are generated by reduction of sulfonyl chlorides. An alternative route is the reaction of Grignard reagents with sulfur dioxide. Transition metal sulfinates are also generated by insertion of sulfur dioxide into metal alkyls, a reaction that may proceed via a metal sulfur dioxide complex. Unsubstituted sulfinic acid, when R is the hydrogen atom, is a higher energy isomer of sulfoxylic acid, both of which are unstable.Sulfite
Sulfites or sulphites are compounds that contain the sulfite ion (or the sulfate(IV) ion, from its correct systematic name), SO2−3. The sulfite ion is the conjugate base of bisulfite. Although its acid (sulfurous acid) is elusive, its salts are widely used.
Sulfites are substances that naturally occur in some foods and the human body. They are also used as regulated food additives.Sulfurous acid
Sulfurous acid (also sulphurous acid) is the chemical compound with the formula H2SO3. There is no evidence that sulfurous acid exists in solution, but the molecule has been detected in the gas phase. The conjugate bases of this elusive acid are, however, common anions, bisulfite (or hydrogen sulfite) and sulfite. Sulfurous acid is an intermediate species in the formation of acid rain from sulfur dioxide.Raman spectra of solutions of sulfur dioxide in water show only signals due to the SO2 molecule and the bisulfite ion, HSO−3. The intensities of the signals are consistent with the following equilibrium:
SO2 + H2O ⇌ HSO−3 + H+ Ka = 1.54×10−2; pKa = 1.81.17O NMR spectroscopy provided evidence that solutions of sulfurous acid and protonated sulfites contains a mixture of isomers, which is in equilibrium:
[H–OSO2]− ⇌ [H–SO3]−When trying to concentrate the solution by evaporation to produce waterless sulfurous acid it will decompose (reversing the forming reaction). In cooling down a clathrate SO2·5 3⁄4H2O will crystallise which decomposes again at 7 °C. Thus sulfurous acid H2SO3 cannot be isolated.Sulfuryl fluoride
Sulfuryl fluoride (also spelled sulphuryl fluoride) is an inorganic compound with the formula SO2F2. It is an easily condensed gas and has properties more similar to sulfur hexafluoride than sulfuryl chloride, being resistant to hydrolysis even up to 150 °C. It is neurotoxic and a potent greenhouse gas, but is widely used as a fumigant insecticide to control termites.Thomagata Patera
Thomagata Patera is a volcano on Jupiter's moon Io. It is located on Io's anti-Jupiter hemisphere at 25.67°N 165.94°W / 25.67; -165.94, to the east of the nearby active volcanoes Volund and Zamama. Thomagata is a kidney-shaped Ionian patera, a type of volcanic crater similar to a caldera, 56 kilometers (35 mi) long, 26 km (16 mi) wide, and 1.2–1.6 km (0.7–1.0 mi) deep. The volcano is currently inactive as a thermal hotspot has never been observed at Thomagata and the bright floor of the patera suggests that it is cold enough for sulfur dioxide and sulfur to condense. Thomagata is located near the center of a low, 100 km (62 mi) wide mesa. The edge of the mesa rises 200 meters (660 ft) above the surrounding plains, however the slope up to the edge of Thomagata Patera is unknown. If the floor of the patera is at the same level as the surrounding plains, the western slope of the mesa would have a grade of 2°. The morphology of this mesa and the pattern of faded lava flows along its slopes radiating away from Thomagata (at least on its eastern side) suggest that Thomagata Patera and the mesa that surrounds it may be a shield volcano, also called a tholus on Io. The irregular margin of the mesa and the lack of debris at the base of its basal scarp suggest that it was modified by sulfur dioxide sapping.Vog
Vog is a form of air pollution that results when sulfur dioxide and other gases and particles emitted by an erupting volcano react with oxygen and moisture in the presence of sunlight. The word is a portmanteau of the words "volcanic" and "smog". The term is in common use in the Hawaiian islands, where the Kīlauea volcano, on the Island of Hawaiʻi (the "Big Island"), has been erupting continuously since January 3, 1983. Based on June 2008 measurements, Kīlauea emits 2,000—4,000 tons of sulfur dioxide (SO2) every day.Wellman–Lord process
The Wellman–Lord process is a regenerable process to remove sulfur dioxide from flue gas (flue-gas desulfurization) without creating a throwaway sludge product.
In this process, sulfur dioxide from flue gas is absorbed in a sodium sulfite solution in water forming sodium bisulfite; other components of flue gas are not absorbed. After lowering the temperature the bisulfite is converted to the sodium pyrosulfite which precipitates.
Upon heating, the two previously described chemical reactions are reversed, and sodium pyrosulfite is converted to a concentrated stream of sulfur dioxide and sodium sulfite. The sulfur dioxide can be used for further reactions (e.g. the production of sulfuric acid), and the sulfite is reintroduced into the process.
Na2SO3 + SO2 + H2O → 2NaHSO3
2NaHSO3 → Na2S2O5↓ + H2O
Na2S2O5 + H2O → 2NaHSO3
2NaHSO3 → Na2SO3 + SO2 + H2OIn its initial version (Crane Station, Maryland, 1968) the process was based on potassium sulfite, but the economic prognosis was poor. Interest in the process occurred because of the worldwide shortage of sulfur in 1967 and resulting high prices; power-plant flue gas was viewed as an additional source of sulfur to relieve the shortage. the later version used sodium sulfite and was installed (as a demonstration system funded by USEPA) at Mitchell Station, Indiana in 1974. It was coupled with the Allied reduction (by natural gas) process to make elemental sulfur which can be shipped anywhere, for example to a sulfuric acid plant. Additional installations of W-L were made in New Mexico. The process has been offered commercially by Davy Powergas in Lakeland, Florida. Because of side reactions forming thiosulfate (nonregenerable), there is a small makeup requirement in the form of trona (sodium carbonate).Wet scrubber
The term wet scrubber describes a variety of devices that remove pollutants from a furnace flue gas or from other gas streams. In a wet scrubber, the polluted gas stream is brought into contact with the scrubbing liquid, by spraying it with the liquid, by forcing it through a pool of liquid, or by some other contact method, so as to remove the pollutants.Wine fault
A wine fault or defect is an unpleasant characteristic of a wine often resulting from poor winemaking practices or storage conditions, and leading to wine spoilage. Many of the compounds that cause wine faults are already naturally present in wine but at insufficient concentrations to be of issue. In fact, depending on perception, these concentrations may impart positive characters to the wine. However, when the concentration of these compounds greatly exceeds the sensory threshold, they replace or obscure the flavors and aromas that the wine should be expressing (or that the winemaker wants the wine to express). Ultimately the quality of the wine is reduced, making it less appealing and sometimes undrinkable.There are many causes for the perception in wine faults, including poor hygiene at the winery, excessive or insufficient exposure of the wine to oxygen, excessive or insufficient exposure of the wine to sulphur, overextended maceration of the wine either pre- or post-fermentation, faulty fining, filtering and stabilization of the wine, the use of dirty oak barrels, over-extended barrel aging and the use of poor quality corks. Outside of the winery, other factors within the control of the retailer or end user of the wine can contribute to the perception of flaws in the wine. These include poor storage of the wine that exposes it to excessive heat and temperature fluctuations as well as the use of dirty stemware during wine tasting that can introduce materials or aromas to what was previously a clean and fault-free wine.
|Mixed oxidation states|
|+1 oxidation state|
|+2 oxidation state|
|+3 oxidation state|
|+4 oxidation state|
|+5 oxidation state|
|+6 oxidation state|
|+7 oxidation state|
|+8 oxidation state|
Oxides are sorted by oxidation state. Category:Oxides