Sulfur dioxide

Sulfur dioxide (also sulphur dioxide) is the smell of rotten eggs.

SO₂ is produced by acid rain.^[2]

Preparation

Sulfur dioxide can be prepared by burning sulfur:

S₈ + 8 O₂ → 8 SO₂

The combustion of hydrogen sulfide and organosulfur compounds proceeds similarly.

2 H₂S(g) + 3 O₂(g) → 2 H₂O(g) + 2 SO₂(g)

The roasting of sulfide ores such as iron cinnabar (mercury sulfide) also releases SO₂:

4 FeS₂(s) + 11 O₂(g) → 2 Fe₂O₃(s) + 8 SO₂(g)

2 ZnS(s) + 3 O₂(g) → 2 ZnO(s) + 2 SO₂(g)

HgS(s) + O₂(g) → Hg(g) + SO₂(g)

Sulfur dioxide is a by-product in the manufacture of coke and sand in this process:

2 CaSO₄(s) + 2SiO₂(s) + C(s) → 2 CaSiO₃(s) + 2 SO₂(g) + CO₂(g)

Action of hot sulfuric acid on copper turnings produces sulfur dioxide.

Cu(s) + 2H₂SO₄(aq) → CuSO₄(aq) + SO₂(g) + 2H₂O(l)

Structure and bonding

SO₂ is a bent molecule with C_2v symmetry point group. In terms of molecular orbital theory, most of these valence electrons are engaged in S-O bonding.

Reactions

Treatment of basic solutions with sulfur dioxide affords sulfite salts:

SO₂ + 2 NaOH → Na₂SO₃ + H₂O

Featuring sulfur in the +4 oxidation state, sulfur doxide is a reducing agent. It is oxidized by halogens such as chlorine to give the sulfuryl halides:

SO₂ + Cl₂ → SO₂Cl₂

However, on rare occasions, it can also act as an Claus process, sulfur dioxide is reduced by hydrogen sulfide to give elemental sulfur:

SO₂ + 2 H₂S → 3 S + 2 H₂O

Uses

As a preservative

Sulfur dioxide is sometimes used as a preservative for dried apricots and other dried fruits due to its antimicrobial properties. The preservative is used to maintain the appearance of the fruit rather than to prevent rotting. Its presence can give fruit a distinctive chemical taste.

In winemaking

Sulfur dioxide is a very important compound in winemaking, and is designated as parts per million in wine, E number: E220.^[3] It is present even in so-called unsulphurated wine at concentrations of up to 10 milligrams per litre.^[4] It serves as an antibiotic and antioxidant, protecting wine from spoilage by bacteria and oxidation. It also helps to keep volatile acidity at desirable levels. Sulfur dioxide is responsible for the words "contains sulfites" found on wine labels. Wines with SO₂ concentrations below 10ppm do not require "contains sulfites" on the label by US and EU laws. The upper limit of SO₂ allowed in wine is 350ppm in US, in the EU is 160 ppm for red wines and 210 ppm for white and rosé wines. In low concentrations SO₂ is mostly undetectable in wine, but at over 50ppm, SO₂ becomes evident in the nose and taste of wine.

SO₂ is also a very important element in winery sanitation. Wineries and equipment must be kept very clean, and because bleach cannot be used in a winery, a mixture of SO₂, water, and citric acid is commonly used to clean hoses, tanks, and other equipment to keep it clean and free of bacteria.

As a reducing bleach

Sulfur dioxide is also a good Oxygen in the atmosphere reoxidizes the reduced dyes, restoring the color.

Precursor to sulfuric acid

Sulfur dioxide is also used to make sulfuric acid, being converted to contact process.

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. Sulfur dioxide has no role in mammalian biology. Sulfur dioxide blocks nerve signals from the pulmonary stretch receptors (PSR's) and abolishes the Hering-Breuer inflation reflex.

As a refrigerant

Being easily condensed and with a high heat of evaporartion, sulfur dioxide is a candidate material for refrigerants. Prior to the development of refrigerant in home refrigerators.

As a reagent and solvent

Sulfur dioxide is a versatile inert solvent that has been widely used for dissolving highly oxidizing salts. It is also used occasionally as a source of the sulfonyl group in organic synthesis. Treatment of aryldiazonium salts with sulfur dioxide affords the corresponding aryl sulfonyl chloride.^[5]

Emissions

According to the U.S. EPA (as presented by the 2002 World Almanac or in chart form^[6]), the following amount of sulfur dioxide was released in the U.S. per year, measured in thousands of short tons:

Due largely to the US EPA’s Acid Rain Program, the U.S. has witnessed a 33 percent decrease in emissions between 1983 and 2002. This improvement resulted from calcium oxide (lime) reacts with sulfur dioxide to form calcium sulfite:

CaO + SO₂ → CaSO₃

Aerobic oxidation converts this CaSO₃ into CaSO₄, gypsum. Most gypsum sold in Europe comes from flue gas desulfurization.

New fuel additive catalysts, such as ferox, are being used in gasoline and diesel engines in order to lower the emission of sulfur oxide gases into the atmosphere. This is also done by forcing the sulfur into stable mineral salts and mixed mineral sulfates as opposed to sulfuric acid and sulfur oxides.

As of 2006, China is the world's largest sulfur dioxide polluter, with 2005 emissions estimated to be 25.49 million tons. This amount represents a 27% increase since 2000, and is roughly comparable with U.S. emissions in 1980^[7].

Al-Mishraq, an Iraqi sulfur plant, was the site of a 2003 disaster resulting in the release of massive amounts of sulfur dioxide into the atmosphere.

Temperature dependence of aqueous solubility

• The values are tabulated for 101.3 kPa Henry's law.
• The solublity is given for "pure water", i.e., water that contains only SO₂ in the amount at equilibrium with the gas phase. This "pure water" is going to be acidic. The solublity of SO₂ in neutral (or alkaline) water is generally going to be higher because of the pH-dependent speciation of SO₂ in the solution with the production of bisulfite and some sulfite ions.

Safety

According to Claude Ribbe in The Crime of Napoleon, sulfur dioxide gas was used as an execution poison by the French emperor to suppress a slave revolt in Haiti early in the 19th century.

References

1. ^ http://chemed.chem.purdue.edu/genchem/topicreview/bp/ch8/table.php
2. ^ Holleman, A. F.; Wiberg, E. "Inorganic Chemistry" Academic Press: San Diego, 2001. ISBN 0-12-352651-5.
3. ^ Current EU approved additives and their E Numbers, The Food Standards Agency website.
4. ^ Sulphites in wine, MoreThanOrganic.com.
5. ^ R. V. Hoffman “m-Trifluoromethylbenzenesulfonyl chloride” Organic Syntheses, Collected Volume 7, p.508 (1990). http://www.orgsyn.org/orgsyn/pdfs/CV7P0508.pdf.
6. ^ National Trends in Sulfur Dioxide Levels, United States Environmental Protection Agency.
7. ^ China has its worst spell of acid rain, United Press International.

See also

• Sulfur
• Sulfur trioxide
• Sulfur-iodine cycle
• National Ambient Air Quality Standards
• Homer City Generating Station