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SULFUR DIOXIDE

Linear Formula: SO2
CAS Number: 7446-09-5
Molecular Weight: 64.06
Beilstein: 3535237
EC Number: 231-195-2
MDL number: MFCD00011450
PubChem Substance ID: 24857804


APPLICATIONS

The overarching, dominant use of sulfur dioxide is in the production of sulfuric acid.
Sulfur dioxide is primarily used as a feedstock for the production of sulfuric acid (contact process). 
Sulfur dioxide is generally used in the synthesis of a variety of sulfur-containing organic compounds like alkyl/aryl sulfonyl chlorides, sulfinates, sultines, and polysulfone. 

Additionally, Sulfur dioxide also possess antimicrobial properties and are used as preservatives.
Sulfur dioxide is a sulfur oxide. 
Sulfur dioxide has a role as a food bleaching agent, a refrigerant and an Escherichia coli metabolite.

Sulfur dioxide is used in many industries. 
Some examples Sulfur dioxide usage in industry are manufacture sulfuric acid, paper, and food preservatives. 

-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 ability to prevent oxidation, 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.

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

Sulfur dioxide is still an important compound in winemaking, and is measured in parts per million (ppm) in wine. 
Sulfur dioxide is present even in so-called unsulfurated wine at concentrations of up to 10 mg/L. 

Sulfur dioxide 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 Sulfur dioxide. 
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) Sulfur dioxide, which is the active form, while at higher pH more Sulfur dioxide is found in the inactive sulfite and bisulfite forms. 

The molecular Sulfur dioxide 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 Sulfur dioxide concentrations below 10 ppm do not require "contains sulfites" on the label by US and EU laws. 
The upper limit of total Sulfur dioxide 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, Sulfur dioxide is mostly undetectable in wine, but at free SO2 concentrations over 50 ppm, SO2 becomes evident in the smell and taste of wine.

Sulfur dioxide 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 cork taint, a mixture of Sulfur dioxide, 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.

-As a reducing agent

Sulfur dioxide is also a good reductant. 
In the presence of water, sulfur dioxide is able to decolorize substances. 
Specifically, Sulfur dioxide 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:

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 Yersinia pestis 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 Sulfurozador, 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 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, and 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. 
Sulfur dioxide lowers blood pressure and slows hypertensive remodeling of blood vessels, especially thickening of their intima. 
Sulfur dioxide also regulates lipid metabolism.

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

-Aspirational applications

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 chlorofluorocarbons, sulfur dioxide was used as a refrigerant in home refrigerators.

-Climate engineering
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 1991 eruption of Mount Pinatubo. 
However this form of geoengineering would have uncertain regional consequences on rainfall patterns, for example in monsoon regions.


DESCRIPTION

Sulfur dioxide (IUPAC-recommended spelling) or sulphur dioxide (traditional Commonwealth English) is the chemical compound with the formula SO2. 
Sulfur dioxide is a toxic gas responsible for the smell of burnt matches. 

Sulfur dioxide is released naturally by volcanic activity and is produced as a by-product of copper extraction and the burning of sulfur-bearing fossil fuels. 
Sulfur dioxide has a pungent smell like nitric acid.
Sulfur dioxide, 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.

Sulfur dioxide appears as a colorless gas with a choking or suffocating odor. 
Boiling point of Sulfur dioxide is -10°C. 
Sulfur dioxide is Heavier than air. 

Sulfur dioxide is Very toxic by inhalation and may irritate the eyes and mucous membranes. 
Under prolonged exposure to fire or heat the Sulfur dioxide containers may rupture violently and rocket. 
Sulfur dioxide is Used to manufacture chemicals, in paper pulping, in metal and food processing. 

Sulfur dioxide is a colorless gas with a pungent odor. 
And Sulfur dioxide is a liquid when under pressure, and it dissolves in water very easily. 
Sulfur dioxide in the air comes mainly from activities such as the burning of coal and oil at power plants or from copper smelting. 
In nature, sulfur dioxide can be released to the air from volcanic eruptions.

Sulfur dioxide, (SO2), inorganic compound, a heavy, colourless, poisonous gas. 
And, Sulfur dioxide is produced in huge quantities in intermediate steps of sulfuric acid manufacture.
Sulfur dioxide has a pungent, irritating odour, familiar as the smell of a just-struck match. 

Occurring in nature in volcanic gases and in solution in the waters of some warm springs, sulfur dioxide usually is prepared industrially by the burning in air or oxygen of sulfur or such compounds of sulfur as iron pyrite or copper pyrite. 
Large quantities of sulfur dioxide are formed in the combustion of sulfur-containing fuels. 
In the atmosphere Sulfur dioxide can combine with water vapour to form sulfuric acid, a major component of acid rain; in the second half of the 20th century, measures to control acid rain were widely adopted. 

Sulfur dioxide is a precursor of the trioxide (SO3) used to make sulfuric acid. 
In the laboratory the gas may be prepared by reducing sulfuric acid (H2SO4) to sulfurous acid (H2SO3), which decomposes into water and sulfur dioxide, or by treating sulfites (salts of sulfurous acid) with strong acids, such as hydrochloric acid, again forming sulfurous acid.

Sulfur dioxide can be liquefied under moderate pressures at room temperatures; the liquid freezes at −73° C (−99.4° F) and boils at −10° C (14° F) under atmospheric pressure. 
Although its chief uses are in the preparation of sulfuric acid, sulfur trioxide, and sulfites, sulfur dioxide also is used as a disinfectant, a refrigerant, a reducing agent, a bleach, and a food preservative, especially in dried fruits.

Sulfur dioxide (SO2) is a colorless, reactive gas with a strong odor. 
Sulfur dioxide comes from a variety of natural and anthropogenic sources. 
The primary anthropogenic sources of sulfur dioxide emissions are the burning of high-sulfur coals and heating oils in power plants, followed by industrial boilers and metal smelting. 
Natural causes contribute anywhere from 35–65% of total sulfur dioxide emissions annually and include sources such as volcanoes.

Sulfur dioxide (SO2) is comprised of one atom of sulfur and two atoms of oxygen, and is a gas at ambient temperatures. 
Sulfur dioxide has a pungent, irritating odor. 
SO2, Sulfur dioxide, is a member of a family of chemicals comprised of sulfur and oxygen that are collectively known as sulfur oxides (SOX).

Sulfur dioxide (SO2) is a colourless gas with a sharp, irritating odour. 
SO2, Sulfur dioxide, is produced by burning fossil fuels and by the smelting of mineral ores that contain sulfur.
Erupting volcanoes can be a significant natural source of sulfur dioxide emissions.


OCCURRENCE

Sulfur dioxide is found on Earth and exists in very small concentrations and in the atmosphere at about 1 ppm.

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 sulfuric acid, and is a key component of the planet's global atmospheric sulfur cycle 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[12] 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.


PRODUCTION

Sulfur dioxide is primarily produced for sulfuric acid manufacture. 
In the United States in 1979, 23.6 million metric tons (26,014,547 US short tons) of sulfur dioxide were used in this way, compared with 150 thousand metric tons (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.

-Combustion routes

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, 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 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:

2 H2S + 3 O2 → 2 H2O + 2 SO2
The roasting of sulfide ores such as pyrite, sphalerite, and cinnabar (mercury sulfide) also releases Sulfur dioxide:

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 tons of Sulfur dioxide.

-Reduction of higher oxides

Sulfur dioxide can also be a byproduct in the manufacture of calcium silicate cement; CaSO4 is heated with coke and sand in this process:

2 CaSO4 + 2 SiO2 + C → 2 CaSiO3 + 2 SO2 + CO2
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.

Cu + 2 H2SO4 → CuSO4 + SO2 + 2 H2O

-From sulfites

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

SO2 + 2 NaOH → Na2SO3 + H2O
The reverse reaction occurs upon acidification:

H+ + HSO3− → SO2 + H2O


AIR POLLUTION

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 Sulfur dioxide 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.

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.

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.

An analysis found that 18 coal-fired power stations in the western Balkans emitted two-and-half times more sulphur dioxide than all 221 coal plants in the EU combined.

Fuel additives such as calcium additives and magnesium carboxylate may be used in marine engines to lower the emission of sulfur dioxide gases into the atmosphere.

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.

SAFETY

-Inhalation: Sulfur dioxide is VERY TOXIC, can cause death. 
Sulfur dioxide Can cause severe irritation of the nose and throat. 

At high concentrations: Sulfur dioxide can cause life-threatening accumulation of fluid in the lungs (pulmonary edema). 
Symptoms of Sulfur dioxide inhalation may include coughing, shortness of breath, difficult breathing and tightness in the chest. 
A single exposure to a high concentration of Sulfur dioxide can cause a long-lasting condition like asthma. 
If this occurs, many things like other chemicals or cold temperatures can easily irritate the airways. 
Symptoms may include shortness of breath, tightness in the chest and wheezing.

-Skin Contact: Sulfur dioxide is CORROSIVE. 
The gas irritates or burns the skin. 
Permanent scarring can result. 
Direct contact with the liquefied gas can chill or freeze the skin (frostbite). 
Symptoms of mild frostbite include numbness, prickling and itching. 
Symptoms of more severe frostbite include a burning sensation and stiffness. 
The skin may become waxy white or yellow. 
Blistering, tissue death and infection may develop in severe cases.

-Eye Contact: Sulfur dioxide is CORROSIVE. 
The gas irritates or burns the eyes. 
Permanent damage including blindness can result. 
Direct contact with the liquefied gas can freeze the eye. 
Permanent eye damage or blindness can result.

-Ingestion: Not a relevant route of exposure (gas).

-Effects of Long-Term (Chronic) Exposure: May harm the respiratory system. 
Can irritate and inflame the airways.

-Carcinogenicity: Not known to cause cancer.


FIRE HAZARD

Containers may explode in heat of fire or they may rupture and release irritating toxic sulfur dioxide. 
Sulfur dioxide has explosive properties when it comes in contact with sodium hydride; potassium chlorate at elevated temperatures; ethanol; ether; zinc ethylsulfurinate at very cool temperatures (-15C); fluorine; chlorine trifluoride and chlorates. 
Sulfur dioxide will react with water or steam to produce toxic and corrosive fumes. 

When the liquid is heated it may release irritating, toxic sulfur dioxide gas. 
Avoid ammonia, monocesium or monopotassium acetylide; dicesium monoxide; iron (II) oxide; tin oxide; lead (IV) oxide; chromium; manganese; molten sodium, powder aluminum and rubidium. 

Sulfur dioxide has explosive properties when it comes in contact with sodium hydride; potassium chlorate at elevated temperatures; ethanol; ether; zinc ethylsulfurinate at very cool temperatures (-15C); fluorine; chlorine trifluoride and chlorates. 
It will react with water or steam to produce toxic and corrosive fumes. 
Hazardous polymerization may not occur.


PROPERTIES OF SULFUR DIOXIDE

Chemical formula: SO2
Molar mass: 64.066 g mol−1
Appearance: Colorless gas
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)
Solubility in water: 94 g/L
Vapor pressure: 237.2 kPa
Acidity (pKa): 1.81
Basicity (pKb): 12.19
Magnetic susceptibility (χ): −18.2·10−6 cm3/mol
Viscosity: 12.82 μPa·s


SYNONYMS

Sulfurous anhydride
Sulfur(IV) oxide
sulfur dioxide
sulphur dioxide
Sulfurous anhydride
7446-09-5
Sulfurous oxide
Fermenicide powder
Fermenticide liquid 
Fermenicide liquid 
Schwefeldioxyd 
Siarki dwutlenek

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