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Sulfur dioxide (-recommended spelling) or sulphur dioxide (traditional Commonwealth English) is the chemical compound with the formula . It is a responsible for the odor of burnt . It is released naturally by volcanic activity and is produced as a by-product of extraction and the burning of -bearing fossil fuels. It was known to alchemists as "volatile spirit of sulfur" since at least 16th century.

(2024). 9780199652723, Oxford University Press. .

Structure and bonding
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 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 of +4 and a of +1.

Sulfur dioxide is found on Earth and exists in very small concentrations in the atmosphere at about 15 ppb.

On other planets, sulfur dioxide can be found in various concentrations, the most significant being the atmosphere of Venus, where it is the third-most abundant atmospheric gas at 150 ppm. There, it reacts with water to form clouds of , and is a key component of the planet's global atmospheric and contributes to . It has been implicated as a key agent in the warming of early , 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 , is 90% sulfur dioxide

(2024). 9783540346814, Springer-Praxis.
and trace amounts are thought to also exist in the atmosphere of Jupiter. The James Webb Space Telescope has observed the presence of sulfur dioxide on the WASP-39b, where it is formed through in the planet's atmosphere.

As an ice, it is thought to exist in abundance on the —as subliming ice or frost on the trailing hemisphere of Io,

(1985). 9789401088916
and in the crust and mantle of Europa, Ganymede, and Callisto, possibly also in liquid form and readily reacting with water. Europa's Hidden Ice Chemistry – NASA Jet Propulsion Laboratory. (2010-10-04). Retrieved on 2013-09-24.

Sulfur dioxide is primarily produced for manufacture (see , but other processes predated that at least since 16th century). In the in 1979, 23.6 million metric tons (26 million U.S. short tons) of sulfur dioxide were used in this way, compared with 150,000 metric tons (165,347 U.S. short tons) used for other purposes. Most sulfur dioxide is produced by the combustion of elemental . Some sulfur dioxide is also produced by roasting and other ores in air.

Combustion routes
Sulfur dioxide is the product of the burning of or of burning materials that contain sulfur:

+ → , ΔH = −297 kJ/mol

To aid combustion, liquified 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 , 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 and organosulfur compounds proceeds similarly. For example:

+ → +

The roasting of sulfide ores such as , , and (mercury sulfide) also releases SO2:Shriver, Atkins. Inorganic Chemistry, Fifth Edition. W. H. Freeman and Company; New York, 2010; p. 414.

2 + → + 4
+ → +
2 FeS + → + 2

A combination of these reactions is responsible for the largest source of sulfur dioxide, volcanic eruptions. These events can release millions of tons of SO2.

Reduction of higher oxides
Sulfur dioxide can also be a byproduct in the manufacture of ; is heated with coke and sand in this process:

2 + 2 + C → 2 + 2 +

Until the 1970s, commercial quantities of sulfuric acid and cement were produced by this process in , England. Upon being mixed with or , and roasted, the sulfate liberated sulfur dioxide gas, used in sulfuric acid production, the reaction also produced calcium silicate, a precursor in cement production. WHITEHAVEN COAST ARCHAEOLOGICAL SURVEY. (2007)

On a laboratory scale, the action of hot concentrated sulfuric acid on copper produces sulfur dioxide.

Cu + 2 →
Tin also reacts with concentrated sulfuric acid but it produces tin(II) sulfate which can later be pyrolyzed at 360°C into tin dioxide and dry sulfur dioxide.

Sn + →
→ +

From sulfites
The reverse reaction occurs upon acidification:

Sulfites results by the action of aqueous base on sulfur dioxide:

Sulfur dioxide is a mild but useful . It is oxidized by halogens to give the sulfuryl halides, such as sulfuryl chloride:

Sulfur dioxide is the in the , which is conducted on a large scale in . Here, sulfur dioxide is reduced by hydrogen sulfide to give elemental sulfur:

The sequential oxidation of sulfur dioxide followed by its hydration is used in the production of sulfuric acid.

+ + →
Sulfur dioxide dissolves in water to give "", which cannot be isolated and is instead an acidic solution of , and possibly , ions.
Ka = 1.54; p Ka = 1.81

Laboratory reactions
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 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- in a cheletropic reaction to form cyclic . This reaction is exploited on an industrial scale for the synthesis of , which is an important solvent in the petrochemical industry.

Sulfur dioxide can bind to metal ions as a 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 1. As a η1-SO2 (S-bonded planar) ligand sulfur dioxide functions as a Lewis base using the lone pair on S. SO2 functions as a in its η1-SO2 (S-bonded pyramidal) bonding mode with metals and in its 1:1 with Lewis bases such as dimethylacetamide and . When bonding to Lewis bases the of SO2 are EA = 0.51 and EA = 1.56.

The overarching, dominant use of sulfur dioxide is in the production of .

Precursor to sulfuric acid
Sulfur dioxide is an intermediate in the production of sulfuric acid, being converted to , and then to , 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 . Several million tons are produced annually for this purpose.

Food preservative
Sulfur dioxide is sometimes used as a for dried , dried , and other , owing to its properties and ability to prevent , and is called 220 Current EU approved additives and their E Numbers, The Food Standards Agency website. when used in this way in Europe. As a preservative, it maintains the colorful appearance of the fruit and prevents . It is also added to sulfured . Sublimed sulfur is ignited and burned in an enclosed space with the fruits. This is usually done outdoors. Fruits may be sulfured by dipping them into an either , or sodium metabisulfite.

Sulfur dioxide was first used in by the Romans, when they discovered that burning sulfur candles inside empty wine vessels keeps them fresh and free from vinegar smell.

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. Sulphites in wine, It serves as an and , 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.Jackson, R.S. (2008) Wine science: principles and applications, Amsterdam; Boston: Elsevier/Academic Press Its antimicrobial action also helps minimize volatile acidity. Wines containing sulfur dioxide are typically labeled with "containing ".

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 , 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 to the risk of , Chlorine Use in the Winery. Purdue University a mixture of SO2, water, and citric acid is commonly used to clean and sanitize equipment. (O3) is now used extensively for sanitizing in wineries due to its efficacy, and because it does not affect the wine or most equipment. Use of ozone for winery and environmental sanitation , Practical Winery & Vineyard Journal.

As a reducing agent
Sulfur dioxide is also a good . In the presence of water, sulfur dioxide is able to decolorize substances. Specifically, it is a useful reducing for and delicate materials such as clothes. This bleaching effect normally does not last very long. 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 .
(1979). 007041677X, McGraw Hill. 007041677X

Sulfur dioxide is fairly soluble in water, and by both IR and Raman spectroscopy; the hypothetical , 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+

As a fumigant
In the beginning of the 20th century, sulfur dioxide was used in Buenos Aires as a fumigant to kill rats that carried the bacterium, which causes bubonic plague. The application was successful, and the application of this method was extended to other areas in South America. In Buenos Aires, where these apparatuses were known as , but later also in Rio de Janeiro, New Orleans and San Francisco, the sulfur dioxide treatment machines were brought into the streets to enable extensive disinfection campaigns, with effective results.

Biochemical and biomedical roles
Sulfur dioxide 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 is considered that endogenous sulfur dioxide plays a significant physiological role in regulating and function, and aberrant or deficient sulfur dioxide metabolism can contribute to several different cardiovascular diseases, such as arterial hypertension, , pulmonary arterial hypertension, and .

It was shown that in children with pulmonary arterial hypertension due to congenital heart diseases the level of 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 cells in blood vessels, via lowering the activity and activating and protein kinase A. Smooth muscle cell proliferation is one of important mechanisms of hypertensive remodeling of blood vessels and their , so it is an important pathogenetic mechanism in arterial hypertension and atherosclerosis.

Endogenous sulfur dioxide in low concentrations causes endothelium-dependent . 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 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.

Endogenous sulfur dioxide also diminishes myocardial damage, caused by hyperstimulation, and strengthens the myocardial antioxidant defense reserve.

As a reagent and solvent in the laboratory
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 with sulfur dioxide and yields the corresponding aryl sulfonyl chloride, for example:

As a result of its very low , it is often used as a low-temperature solvent/diluent for superacids like (FSO3H/SbF5), allowing for highly reactive species like tert-butyl cation to be observed spectroscopically at low temperature (though tertiary carbocations do react with SO2 above about –30 °C, and even less reactive solvents like SO2ClF must be used at these higher temperatures).

As a refrigerant
Being easily condensed and possessing a high heat of evaporation, sulfur dioxide is a candidate material for refrigerants. Before the development of chlorofluorocarbons, sulfur dioxide was used as a in home refrigerators.


In the United States, the Center for Science in the Public Interest lists the two food preservatives, sulfur dioxide and , as being safe for human consumption except for certain 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. Sulphites may also cause symptoms in non-asthmatic individuals, namely , , flushing, , and , and even life-threatening .

Incidental exposure to sulfur dioxide is routine, e.g. the smoke from , , and sulfur-containing fuels like . Relative to other chemicals, it is only mildly toxic and requires high concentrations to be actively hazardous. Sulfur Dioxide Basics U.S. Environmental Protection Agency However, its ubiquity makes it a major air pollutant with significant impacts on human health. Sulfur Dioxide (SO2) Pollution. United States Environmental Protection Agency

In 2008, the American Conference of Governmental Industrial Hygienists reduced the short-term exposure limit to 0.25 parts per million (ppm). In the US, the OSHA set the PEL at 5 ppm (13 mg/m3) time-weighted average. Also in the US, set the at 100 ppm. In 2010, the "revised the primary SO2 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."

Environmental role

Air pollution
Major volcanic eruptions have an overwhelming effect on sulfate aerosol concentrations in the years when they occur: eruptions ranking 4 or greater on the Volcanic Explosivity Index inject and water vapor directly into the , where they react to create sulfate aerosol plumes. Volcanic emissions vary significantly in composition, and have complex chemistry due to the presence of ash particulates and a wide variety of other elements in the plume. Only containing primarily magmas are responsible for these fluxes, as magma erupted in doesn't result in plumes which reach the stratosphere. However, before the Industrial Revolution, dimethyl sulfide pathway was the largest contributor to sulfate aerosol concentrations in a more average year with no major volcanic activity. According to the IPCC First Assessment Report, published in 1990, volcanic emissions usually amounted to around 10 million tons in 1980s, while dimethyl sulfide amounted to 40 million tons. Yet, by that point, the global human-caused emissions of sulfur into the atmosphere became "at least as large" as all natural emissions of sulfur-containing compounds combined: they were at less than 3 million tons per year in 1860, and then they increased to 15 million tons in 1900, 40 million tons in 1940 and about 80 millions in 1980. The same report noted that "in the industrialized regions of and , anthropogenic emissions dominate over natural emissions by about a factor of ten or even more".IPCC, 1990: Chapter 1: Greenhouse Gases and Aerosols R.T.. In: Climate Change: The IPCC Scientific Assessment J.T.Houghton,. Cambridge University Press, Cambridge, United Kingdom and New York, NY, USA, pp. 31–34, In the eastern United States, sulfate particles were estimated to account for 25% or more of all air pollution. Meanwhile, the Southern Hemisphere had much lower concentrations due to being much less densely populated, with an estimated 90% of the human population in the north. In the early 1990s, anthropogenic sulfur dominated in the Northern Hemisphere, where only 16% of annual sulfur emissions were natural, yet amounted for less than half of the emissions in the Southern Hemisphere. Such an increase in sulfate aerosol emissions had a variety of effects. At the time, the most visible one was , caused by from clouds carrying high concentrations of sulfate aerosols in the . At its peak, acid rain has eliminated and some other fish species and insect life from and in geographically sensitive areas, such as Adirondack Mountains in the . Acid rain worsens function as some of its is lost and heavy metals like are mobilized (spread more easily) while essential nutrients and minerals such as can leach away because of the same. Ultimately, plants unable to tolerate lowered pH are killed, with montane forests being some of the worst-affected due to their regular exposure to sulfate-carrying fog at high altitudes.US EPA: Effects of Acid Rain – Forests While acid rain was too dilute to affect human health directly, breathing or even any air with elevated sulfate concentrations is known to contribute to and conditions, including and . Effects of Acid Rain – Human Health . (June 2, 2006). Retrieved on 2013-02-09. Further, this form of pollution is linked to and low birth weight, with a study of 74,671 pregnant women in finding that every additional 100 µg/m3 of in the air reduced infants' weight by 7.3 g, making it and other forms of air pollution the largest attributable risk factor for low birth weight ever observed.

Control measures
Due largely to the US EPA's Acid Rain Program, the U.S. has had a 33% decrease in emissions between 1983 and 2002 (see table). This improvement resulted in part from flue-gas desulfurization, a technology that enables SO2 to be chemically bound in burning sulfur-containing or .

In particular, reacts with sulfur dioxide to form :

CaO + SO2 → CaSO3

Aerobic oxidation of the CaSO3 gives CaSO4, . Most gypsum sold in Europe comes from flue-gas desulfurization.

To control sulfur emissions, dozens of methods with relatively high efficiencies have been developed for fitting of coal-fired power plants. Sulfur can be removed from coal during burning by using limestone as a bed material in fluidized bed combustion.

(2024). 9781591260493, Professional Publications, Inc.

Sulfur can also be removed from fuels before burning, preventing formation of SO2 when the fuel is burnt. The is used in refineries to produce sulfur as a byproduct. The Stretford process has also been used to remove sulfur from fuel. processes using iron oxides can also be used, for example, Lo-Cat FAQ's About Sulfur Removal and Recovery using the LO-CAT® Hydrogen Sulfide Removal System. or Sulferox. Process screening analysis of alternative gas treating and sulfur removal for gasification. (December 2002) Report by SFA Pacific, Inc. prepared for U.S. Department of Energy (PDF) Retrieved on 2011-10-31.

Fuel additives such as additives and magnesium carboxylate may be used in marine engines to lower the emission of sulfur dioxide gases into the atmosphere.May, Walter R. Marine Emissions Abatement . SFA International, Inc., p. 6.

Impact on climate change

Hydrological cycle

Projected impacts

Solar geoengineering

Table of thermal and physical properties of saturated liquid sulfur dioxide:
(2024). 9780072406559, McGraw-Hill Companies, Inc..
(2024). 9780471457282, John Wiley and Sons, Inc..
Temperature (°C)Density (kg/m^3)Specific heat (kJ/kg K)Kinematic viscosity (m^2/s)Conductivity (W/m K)Thermal diffusivity (m^2/s)Prandtl NumberBulk modulus (K^-1)

See also

External links

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